Organic Electrochemistry: Revised and Expanded [5 ed.] 1420084011, 9781420084016

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FIFTH EDITION

ORGANIC ELECTROCHEMISTRY R E V I S E D A N D E X PA N D E D

© 2016 by Taylor & Francis Group, LLC

FIFTH EDITION

ORGANIC ELECTROCHEMISTRY R E V I S E D A N D E X PA N D E D EDITED BY

Ole Hammerich Uni ve rs i ty o f Co pe nha ge n, Co pe nha ge n, De nma rk

Bernd Speiser Uni ve rs i tä t Tübi nge n, Tübi nge n, Ge rma ny

Boca Raton London New York

CRC Press is an imprint of the Taylor & Francis Group, an informa business

© 2016 by Taylor & Francis Group, LLC

CRC Press Taylor & Francis Group 6000 Broken Sound Parkway NW, Suite 300 Boca Raton, FL 33487-2742 © 2016 by Taylor & Francis Group, LLC CRC Press is an imprint of Taylor & Francis Group, an Informa business No claim to original U.S. Government works Version Date: 20150731 International Standard Book Number-13: 978-1-4200-8402-3 (eBook - PDF) This book contains information obtained from authentic and highly regarded sources. Reasonable efforts have been made to publish reliable data and information, but the author and publisher cannot assume responsibility for the validity of all materials or the consequences of their use. The authors and publishers have attempted to trace the copyright holders of all material reproduced in this publication and apologize to copyright holders if permission to publish in this form has not been obtained. If any copyright material has not been acknowledged please write and let us know so we may rectify in any future reprint. Except as permitted under U.S. Copyright Law, no part of this book may be reprinted, reproduced, transmitted, or utilized in any form by any electronic, mechanical, or other means, now known or hereafter invented, including photocopying, microfilming, and recording, or in any information storage or retrieval system, without written permission from the publishers. For permission to photocopy or use material electronically from this work, please access www.copyright.com (http:// www.copyright.com/) or contact the Copyright Clearance Center, Inc. (CCC), 222 Rosewood Drive, Danvers, MA 01923, 978-750-8400. CCC is a not-for-profit organization that provides licenses and registration for a variety of users. For organizations that have been granted a photocopy license by the CCC, a separate system of payment has been arranged. Trademark Notice: Product or corporate names may be trademarks or registered trademarks, and are used only for identification and explanation without intent to infringe. Visit the Taylor & Francis Web site at http://www.taylorandfrancis.com and the CRC Press Web site at http://www.crcpress.com

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Contents Preface...............................................................................................................................................ix Acknowledgments .......................................................................................................................... xiii Editors .............................................................................................................................................. xv Contributors ...................................................................................................................................xvii

SECTION I Fundamentals and Methods Chapter 1

Basic Concepts ............................................................................................................. 3 Christian Amatore

Chapter 2

Techniques for Studies of Electrochemical Reactions in Solution.............................97 Ole Hammerich and Bernd Speiser

Chapter 3

In Situ Spectroelectrochemistry of Organic Compounds ........................................ 169 Peter Rapta, Evgenia Dmitrieva, Alexey A. Popov, and Lothar Dunsch

Chapter 4

Surface Techniques .................................................................................................. 191 Mohamed M. Chehimi and Jean Pinson

Chapter 5

Application of Digital Simulation ............................................................................ 205 Bernd Speiser

Chapter 6

Theoretical Calculation of Reduction Potentials...................................................... 229 Junming Ho, Michelle L. Coote, Christopher J. Cramer, and Donald G. Truhlar

SECTION II General Preparative Aspects Chapter 7

Preparative Electrolysis on the Laboratory Scale .................................................... 263 Jakob Jörissen and Bernd Speiser

Chapter 8

Application of Ionic Liquids, Emulsions, Sonication, and Microwave Assistance ..... 331 John D. Watkins and Frank Marken

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Chapter 9

Contents

Combinatorial Electrochemistry and Miniaturization ............................................. 345 Kevin D. Moeller

Chapter 10 Relations between Micro- and Macrophenomena.................................................... 371 Christian Amatore

SECTION III

Electron Transfers and Concerted Processes

Chapter 11 Influence of Molecular and Medium Effects on Two-Electron Processes ............... 395 Kevin Lam and William E. Geiger Chapter 12 Electrochemically Driven Supramolecular Devices ................................................ 433 Paola Ceroni, Alberto Credi, and Margherita Venturi Chapter 13 Proton-Coupled Electron Transfers .......................................................................... 481 Cyrille Costentin, Marc Robert, and Jean-Michel Savéant Chapter 14 Dissociative Electron Transfers ................................................................................ 511 Cyrille Costentin, Marc Robert, and Jean-Michel Savéant Chapter 15 Electron Transfer–Catalyzed Reactions ................................................................... 531 Kazuhiro Chiba and Yohei Okada

SECTION IV

Organic Electrochemical Reaction Types

Chapter 16 Cleavages and Deprotections ................................................................................... 559 Ole Hammerich Chapter 17 Reductive Coupling .................................................................................................. 621 James H.P. Utley, R. Daniel Little, and Merete Folmer Nielsen Chapter 18 Oxidative Coupling .................................................................................................. 705 Hans J. Schäfer Chapter 19 Oxidative Substitution and Addition Reactions ....................................................... 775 Ole Hammerich and James H.P. Utley Chapter 20 Fluorination ..............................................................................................................807 Toshio Fuchigami and Shinsuke Inagi

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SECTION V Electrochemical Conversions of Organic Compounds Chapter 21 Electrochemistry of Fullerenes, Derivatives, and Related Compounds................... 829 Frederic Melin, Lourdes E. Echegoyen, and Luis Echegoyen Chapter 22 Aliphatic and Aromatic Hydrocarbons: Reduction .................................................. 861 Jürgen Heinze Chapter 23 Oxidation of Hydrocarbons ...................................................................................... 891 Ole Hammerich Chapter 24 Activation of the Carbon–Halogen Bond ................................................................. 917 Armando Gennaro, Abdirisak Ahmed Isse, and Patrizia Romana Mussini Chapter 25 Aliphatic and Aromatic Halides: Conversions ......................................................... 941 Dennis G. Peters Chapter 26 Oxygen-Containing Compounds: Alcohols, Ethers, and Phenols ........................... 981 Robert Francke, Thomas Quell, Anton Wiebe, and Siegfried R. Waldvogel Chapter 27 Sulfur-, Selenium-, and Tellurium-Containing Compounds .................................. 1035 Richard S. Glass Chapter 28 Aliphatic Nitrogen–Containing Compounds: Amines, Amino Alcohols, and Amino Acids.................................................................................................... 1103 Osamu Onomura Chapter 29 Aromatic Nitrogen–Containing Compounds ..........................................................1121 Jan S. Jaworski Chapter 30 Reduction of Nitro Compounds and Related Substrates .........................................1149 Ole Hammerich Chapter 31 Reduction of Aldehydes, Ketones, and Azomethines ............................................ 1201 Jiří Ludvík Chapter 32 Reductions of Carboxylic Acids and Derivatives................................................... 1249 Rolf Breinbauer and Martin Peters Chapter 33 Oxidation of Carboxylic Acids and Derivatives .................................................... 1267 Hideo Tanaka, Manabu Kuroboshi, and Sigeru Torii

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Chapter 34 Heterocyclic Compounds ....................................................................................... 1309 Fructuoso Barba and Belen Batanero Chapter 35 Organoelemental Compounds ................................................................................ 1357 Jun-ichi Yoshida, Toshiki Nokami, and Seiji Suga Chapter 36 Organometallic Compounds as Tools in Organic Electrosynthesis ....................... 1393 Anny Jutand

SECTION VI

Stereochemical and Biological Aspects

Chapter 37 Electrosynthesis of Bioactive Materials ................................................................. 1435 Randi K. Gbur and R. Daniel Little Chapter 38 Stereochemistry of Organic Electrode Processes .................................................. 1461 Toshio Fuchigami and Shinsuke Inagi Chapter 39 Electroenzymatic Synthesis ....................................................................................1511 Christina Kohlmann and Stephan Lütz Chapter 40 Electrochemical Modeling of Biological Processes ............................................... 1543 Richard D. Webster

SECTION VII

Surface Confined Systems

Chapter 41 Electrochemistry of Conducting Polymers ............................................................ 1571 Jürgen Heinze Chapter 42 Surface-Bound and Immobilized Molecules ......................................................... 1605 Jean Pinson

SECTION VIII Special Applications Chapter 43 Electrogenerated Bases and Nucleophiles.............................................................. 1625 James H.P. Utley, Merete Folmer Nielsen, and Peter B. Wyatt Chapter 44 Electrocatalytic Hydrogenation .............................................................................. 1657 Jean Lessard Index ............................................................................................................................................ 1673

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Preface Organic electrochemistry is concerned with the reduction and oxidation of organic molecules at electrodes. Although it is now more than 200 years ago that the so-called Volta pile was discovered, it was not until 1830–1850 that investigations of organic electrochemical processes, pioneered by Faraday and Kolbe, were established as a research area in their own right. Toward the end of the  nineteenth century, investigators such as Tafel and Haber made significant contributions to the present knowledge of organic electrode processes. Haber, for instance, in his now famous paper on the reduction of nitrobenzene (F. Haber, Z. Elektrochem. 1898, 4, 506) recognized the significance of the electrode potential in the following words: “Oxydations- und Reduktionsprozesse hängen in erster Linie von dem Potential der Elektrode ab, an welcher sie ablaufen, und Stromdichte, Stromdauer und Elektrodenmaterial sind nur insofern bedeutsam, als sie das Elektrodenpotential und seine Änderungen im Gang der Elektrolyse bestimmen.” (“Oxidation and reduction processes primarily depend on the potential of the electrode at which they proceed, and current density, current duration, and electrode material are only important insofar as they determine the electrode potential and its changes during the electrolysis.”) The application of electrolysis as a means of preparing organic compounds continued in the first half of the twentieth century. This development took place along with the development of new electrochemical techniques for the study of electrode processes, for instance, polarography at the dropping mercury electrode introduced by Heyrovsky in the early 1920s. Other important contributions were due to Lingane, Kolthoff, Laitinen, and Delahay. Later, Hickling’s potentiostat (A. Hickling, Trans. Faraday Soc. 1942, 38, 27) as an experimental tool to control experiments led into the computerization era. Most of the work reported before World War II was carried out in aqueous electrolyte solutions. This situation changed after the war, and since the mid-1950s, the attention has been focused mostly on the application of nonaqueous solvents. This has allowed for the detection of the primary intermediates, typically radical anions and radical cations, and for the study of their reactions. The theoretical foundations for the analysis of kinetics and mechanisms by, for instance, cyclic voltammetry and related techniques were mostly published in the 1960s and 1970s. The application of such techniques has resulted in a steadily increasing understanding of the kinetics and mechanisms of organic electrochemical processes. Facing the fact that thousands of organic electrochemical processes are now known, it is striking that most often the only electrochemical reaction to be mentioned in a typical organic chemistry textbook is the oxidation of an organic carboxylic acid R-COOH to the corresponding dimeric alkane R-R reported by Kolbe as early as 1849. Also, it is often overlooked that reductions by metals involving radical anions as intermediates are intimately related to cathodic reductions. And most frequently, radical cations are not mentioned at all, except, of course, in the context of mass spectrometry! This problem, that organic electrochemistry has had difficulties in penetrating into the organic chemistry curriculum, was pointed out already in the preface to the first edition of this book. Unfortunately, not much has changed in the more than 40 years that have elapsed between that edition and this fifth edition. The knowledge of organic electrochemistry had in the 1960s and the early 1970s matured to the point where the time was ripe for the first edition of Organic Electrochemistry (M.M. Baizer, ed., Organic Electrochemistry, Marcel Dekker: New York, 1973). The editor was Manual M. Baizer, well known for his contributions to preparative organic electrochemistry, first of all the development of the electrohydrodimerization of acrylonitrile into a highly successful industrial process for the manufacture of adiponitrile. Baizer set the standards for the book that was organized to include chapters both on the electrochemical reduction and oxidation of specific classes of compounds and

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Preface

on specific types of electrode processes. This approach obviously led to some overlap, a problem that Baizer touched upon in the following words: “It has become routine, at least for reviewers, to point out that in a multi-authored book there is overlap of material, non-uniformity of style, repetition etc. This book was not hastily assembled, and there was adequate time to achieve uniformity if that had been desired. But the material was deliberately organized so that a given segment might appear in two or more contexts and, further, where unanimity of opinion does not exist, more than one current viewpoint might be expressed.” The same basic philosophy was followed in the organization of the second (Baizer, M.M., Lund, H., eds., Organic Electrochemistry, Marcel Dekker: New York, 1983), third (Lund, H., Baizer, M.M., eds., Organic Electrochemistry, Marcel Dekker: New York, 1991), fourth (Lund, H., Hammerich, O., eds., Organic Electrochemistry, Marcel Dekker: New York, 2001) editions, and now for this fifth edition. Owing to this intentional overlap between chapters, the reader is encouraged to use the index to find details, in addition to those found in a given chapter, of the electrochemistry of a specific compound or a class of compounds. Also, some formal variations that reflect common usage by different authors have not been brought into line, for example, the use of “e” or “e −” for the electron or the nomenclature of electrochemical mechanisms (“ECE” vs. “eCe,” etc.). The progress in organic electrochemistry that has taken place since the appearance of the fourth edition is reflected by the organization of this edition. Organic electrochemical reactions are rather complex with electron transfers and transport processes interwoven with bond breaking and forming reaction steps. As a consequence, different perspectives have developed over the years and organic electrochemistry may be discussed following different lines of thought. Moreover, organic electrochemical reactions are often integrated into more complicated synthetic strategies and advanced chemical reasoning or are applied to complex materials and structures. Consequently, it was decided to present the field as follows. After an introduction of some basic and technical aspects, the reaction step that distinguishes organic electrochemistry from classical organic reactions, the electron transfer to and from organic molecules, is described and its most important variants are presented. Second, the importance of electron transfers for the initiation of various organic electrochemical reaction types (follow-up reactions) is emphasized. Third, the electrochemical transformations of organic compounds are systematically presented according to the type of starting materials. Finally, cases where organic electrochemistry forms an integral part in a wider context of chemical research (e.g., in biological or materials science) are discussed. In comparison with the fourth edition, the following major changes have been made to this edition: Sixteen new chapters have been included. These are: • Chapter 3: In Situ Spectroelectrochemistry of Organic Compounds by Peter Rapta, Evgenia Dmitrieva, Alexey A. Popov, and Lothar Dunsch • Chapter 4: Surface Techniques by Mohamed M. Chehimi, and Jean Pinson • Chapter 5: The Application of Digital Simulation by Bernd Speiser • Chapter 6: Theoretical Calculation of Reduction Potentials by Junming Ho, Michelle L. Coote, Christopher J. Cramer, and Donald G. Truhlar • Chapter 8: The Application of Ionic Liquids, Emulsions, Sonication, and Microwave Assistance by John D. Watkins and Frank Marken • Chapter 9: Combinatorial Electrochemistry and Miniaturization by Kevin D. Moeller • Chapter 11: Influence of Molecular and Medium Effects on Two-Electron Processes by Kevin Lam and William E. Geiger • Chapter 12: Electrochemically Driven Supramolecular Devices by Paola Ceroni, Alberto Credi, and Margherita Venturi • Chapter 13 and 14: Proton-Coupled Electron Transfers and Dissociative Electron Transfers by Cyrille Costentin, Marc Robert, and Jean-Michel Savéant • Chapter 15: Electron Transfer Catalyzed Reactions by Kazuhiro Chiba and Yohei Okada

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Preface

• Chapter 24: Activation of the Carbon-Halogen Bond by Armando Gennaro, Abdirisak Ahmed Isse, and Patrizia Romana Mussini • Chapter 36: Organometallic Compounds as Tools in Organic Electrosynthesis by Anny Jutand • Chapter 40: Electrochemical Modeling of Biological Processes by Richard D. Webster • Chapter 42: Surface-Bound and Immobilized Molecules by Jean Pinson • Chapter 44: Electrocatalytic Hydrogenation by Jean Lessard Approximately one-fourth of the previous chapters have been rewritten by new authors: • Chapter 7: Preparative Electrolysis on the Laboratory Scale by Jakob Jörissen and Bernd Speiser • Chapter 16: Cleavages and Deprotections by Ole Hammerich • Chapter 26: Oxygen Containing Compounds: Alcohols, Ethers, and Phenols by Robert Francke, Thomas Quell, Anton Wiebe, and Siegfried R. Waldvogel • Chapter 27: Sulfur, Selenium, and Tellurium Containing Compounds by Richard S. Glass • Chapter 30: Reduction of Nitro Compounds and Related Substrates by Ole Hammerich • Chapter 31: Reduction of Aldehydes, Ketones, and Azomethines by Jiri Ludvik • Chapter 32: Reduction of Carboxylic Acids and Derivatives by Rolf Breinbauer and Martin Peters • Chapter 34: Heterocyclic Compounds by Fructuoso Barba and Belen Batanero • Chapter 39: Electroenzymatic Synthesis by Christina Kohlmann and Stephan Lütz All other chapters have been thoroughly revised or updated where needed, in some cases in cooperation with new coauthors. The following four chapters have been omitted: • Old Chapter 4: Comparison Between Electrochemical Reactions and Chemical Oxidations and Reductions • Old Chapter 28: Amalgam and Related Reductions • Old Chapter 29: Electrogenerated Reagents • Old Chapter 31: Industrial Electroorganic Chemistry Thus, in a sense, this is a new book and not just a conservative update to the previous edition. Altogether this book encompasses 44 chapters written by 66 authors. In most cases, we have encouraged authors to avoid back references to the fourth edition. However, this was not always possible, nor desirable. The text of some chapters in the fourth edition is still authoritative and not much could be added (see, e.g., the appendices on solvents and supporting electrolytes in Chapter 7). Many of the authors of the fourth and previous editions have now retired. Fortunately, new authors have willingly accepted to carry the torch on to this new edition. Sadly, one of these new authors, Lothar Dunsch, passed away in late 2013. This was a great loss to the electrochemical community, and he will be remembered as one of the leading figures in the development and application of spectroelectrochemistry. Ole Hammerich Bernd Speiser

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Acknowledgments The editors thank all the authors for the enthusiasm with which they have undertaken the enormous task of making this fifth edition possible. We have strived to maintain the basic philosophy and not least the high standards of the previous editions to make this a worthy fifth edition of the wellknown “Baizer.” Hopefully, we have succeeded. We acknowledge the help of several people who have contributed “behind the scenes” in some way or the other. Klaus-Michael Mangold (Frankfurt/Main, Germany), Susanne Hempel (Tübingen, Germany), and Tatiana V. Magdesieva (Moscow, Russia) helped with locating or citing some unusual references. Practically, each and every manuscript file (often in a multitude of versions) went through the hands (and across the fingertips) of Britta Rochier (Tübingen, Germany), who almost always found a way to cope with the astonishingly variable word processing software in use around the world. Finally, we thank the staff at Taylor & Francis Group, Barbara Glunn, Kari A. Budyk, and Jennifer Derima, for their patience with our inquiries while we were collecting the manuscripts from the authors.

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Editors Ole Hammerich is an associate professor in the Department of Chemistry, University of Copenhagen, Denmark, where he has taught since 1975. Dr. Hammerich’s research interests include kinetics and mechanisms of organic radical ion reactions, charge-transfer complexes, and electropolymerization. He is author or coauthor of more than 110 articles, papers, and book chapters and was, together with Henning Lund, the editor of the fourth edition of Organic Electrochemistry. Together with Jens Ulstrup, he was the editor of Bioinorganic Electrochemistry. Dr. Hammerich has served as the national secretary for Denmark and as the chairman of the molecular electrochemistry division in the International Society of Electrochemistry and he is a member of the Danish Electrochemical Society and the Electrochemical Society. He earned his PhD (1974) from the University of Copenhagen, Denmark. Bernd Speiser is an außerplanmäßiger professor in the Institut für Organische Chemie, Universität Tübingen, Germany. He received his doctoral degree in Tübingen in 1981 and finished his habilitation in 1990. He was a Heisenberg fellow of the Deutsche Forschungsgemeinschaft in the years 1992–1996. He teaches courses in organic chemistry and molecular electrochemistry. His research is focused on the elucidation of reaction mechanisms of organic, organometallic, and inorganic compounds with experimental methods and computer simulation. This includes the study of multielectron transfers, transport properties, redox-active nanoparticles, and combinatorial electrochemistry. He has published more than 140 full papers, communications, and book chapters. He was the founding chair of Division 6 (molecular electrochemistry) in the International Society of Electrochemistry and is a member of the Electrochemical Society, the International Society of Electrochemistry, the Society of Electroanalytical Chemistry, the Gesellschaft Deutscher Chemiker (GDCh), and DECHEMA Gesellschaft für chemische Technik und Biotechnologie.

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Contributors Christian Amatore Ecole Normale Supérieure Department of Chemistry (ENS-CNRS-UPMC) PSL Research Univerisity Sorbonne Universités Paris, France

Michelle L. Coote ARC Centre of Excellence for Electromaterials Science Research School of Chemistry Australian National University Canberra, Australian Capital Territory, Australia

Fructuoso Barba Department of Organic Chemistry University of Alcala Madrid, Spain

Cyrille Costentin Laboratoire d’Electrochimie Moléculaire, UMR CNRS 7591 Université Paris Diderot Sorbonne Paris Cité Paris, France

Belen Batanero Department of Organic Chemistry University of Alcala Madrid, Spain Rolf Breinbauer Institute of Organic Chemistry Graz University of Technology Graz, Austria Paola Ceroni Dipartimento di Chimica “G. Ciamician” Alma Mater Studiorum Università di Bologna Bologna, Italy Mohamed M. Chehimi Université Paris Diderot Sorbonne Paris Cité, ITODYS, UMR 7086 CNRS Paris, France Kazuhiro Chiba Department of Applied Biological Chemistry Tokyo University of Agriculture and Technology Tokyo, Japan

Christopher J. Cramer Department of Chemistry, Chemical Theory Center, and Supercomputing Institute University of Minnesota Minneapolis, Minnesota Alberto Credi Dipartimento di Chimica “G. Ciamician” Alma Mater Studiorum Università di Bologna Bologna, Italy Evgenia Dmitrieva Department of Electrochemistry and Conducting Polymers Center of Spectroelectrochemistry Leibniz Institute for Solid State and Materials Research Dresden, Germany Lothar Dunsch (deceased) Department of Electrochemistry and Conducting Polymers Center of Spectroelectrochemistry Leibniz Institute for Solid State and Materials Research Dresden, Germany

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Lourdes E. Echegoyen Department of Chemistry University of Texas at El Paso El Paso, Texas Luis Echegoyen Department of Chemistry University of Texas at El Paso El Paso, Texas Robert Francke Department of Chemistry University of Rostock Rostock, Germany Toshio Fuchigami Department of Electronic Chemistry Tokyo Institute of Technology Yokohama, Japan Randi K. Gbur Department of Chemistry and Biochemistry University of California, Santa Barbara Santa Barbara, California

Contributors

Junming Ho Institute of High Performance Computing Singapore Shinsuke Inagi Department of Electronic Chemistry Tokyo Institute of Technology Yokohama, Japan Abdirisak Ahmed Isse Dipartimento di Scienze Chimiche Università degli Studi di Padova Padova, Italy Jan S. Jaworski Faculty of Chemistry University of Warsaw Warszawa, Poland Jakob Jörissen Lehrstuhl für Technische Chemie (Chemische Prozessentwicklung) Technische Universität Dortmund Dortmund, Germany

William E. Geiger University of Vermont Burlington, Vermont

Anny Jutand Département de Chimie Ecole Normale Supérieure Paris, France

Armando Gennaro Dipartimento di Scienze Chimiche Università degli Studi di Padova Padova, Italy

Christina Kohlmann BASF Personal Care and Nutrition GmbH Düsseldorf, Germany

Richard S. Glass Department of Chemistry and Biochemistry The University of Arizona Tucson, Arizona Ole Hammerich Department of Chemistry University of Copenhagen Copenhagen, Denmark Jürgen Heinze Institut für Physikalische Chemie Freiburger Materialforschungszentrum Universität Freiburg Freiburg, Germany

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Manabu Kuroboshi School of Natural Science and Technology Okayama University Okayama, Japan Kevin Lam School of Science and Technology Nazarbayev University Astana, Kazakhstan Jean Lessard Laboratoire de Chimie et Électrochmie Organiques Département de Chimie Université de Sherbrooke Sherbrooke, Québec, Canada

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R. Daniel Little Department of Chemistry and Biochemistry University of California, Santa Barbara Santa Barbara, California

Osamu Onomura Nagasaki University Nagasaki, Japan

Jiří Ludvík Department of Molecular Electrochemistry J. Heyrovský Institute of Physical Chemistry Academy of Sciences of the Czech Republic Prague, Czech Republic

Dennis G. Peters Department of Chemistry Indiana University Bloomington, Indiana

Stephan Lütz Global Discovery Chemistry Novartis Pharma AG Basel, Switzerland Frank Marken Department of Chemistry University of Bath Bath, United Kingdom Frederic Melin Laboratoire de Bioélectrochimie et Spectroscopie Université de Strasbourg Strasbourg, France Kevin D. Moeller Department of Chemistry Washington University in St. Louis St. Louis, Missouri Patrizia Romana Mussini Dipartimento di Chimica Università degli Studi di Milano  Milan, Italy Merete Folmer Nielsen The Danish Emergency Management Agency Birkerød, Denmark Toshiki Nokami Department of Chemistry and Biotechnology Tottori University Tottori, Japan Yohei Okada Graduate School of Biomedical Sciences Tokyo University of Agriculture and Technology Tokyo, Japan

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Martin Peters Institute of Organic Chemistry Graz University of Technology Graz, Austria Jean Pinson Université Paris Diderot Sorbonne Paris Cité, ITODYS, UMR 7086 CNRS Paris, France Alexey A. Popov Department of Electrochemistry and Conducting Polymers Center of Spectroelectrochemistry Leibniz Institute for Solid State and Materials Research Dresden, Germany Thomas Quell Department of Organic Chemistry Johannes Gutenberg University Mainz, Germany Peter Rapta Institute of Physical Chemistry and Chemical Physics Slovak University of Technology Bratislava Bratislava, Slovak Republic Marc Robert Laboratoire d’Electrochimie Moléculaire, UMR CNRS 7591 Université Paris Diderot Sorbonne Paris Cité Paris, France

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Jean-Michel Savéant Laboratoire d’Electrochimie Moléculaire, UMR CNRS 7591 Université Paris Diderot Sorbonne Paris Cité Paris, France Hans J. Schäfer Organisch-Chemisches Institut Westfälische Wilhelms-Universität Münster, Germany Bernd Speiser Institut für Organische Chemie Universität Tübingen Tübingen, Germany Seiji Suga Division of Applied Chemistry Graduate School of Natural Science and Technology Okayama University Okayama, Japan Hideo Tanaka School of Natural Science and Technology Okayama University Okayama, Japan Sigeru Torii The Institute of Creative Chemistry Co., Ltd. Okayama-City, Japan Donald G. Truhlar Department of Chemistry, Chemical Theory Center, and Supercomputing Institute University of Minnesota Minneapolis, Minnesota James H.P. Utley School of Biological and Chemical Sciences Queen Mary University of London London, United Kingdom

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Contributors

Margherita Venturi Dipartimento di Chimica “G. Ciamician” Alma Mater Studiorum Università di Bologna Bologna, Italy Siegfried R. Waldvogel Department of Organic Chemistry Johannes Gutenberg University Mainz, Germany John D. Watkins Department of Chemistry University of Bath Bath, United Kingdom Richard D. Webster Division of Chemistry and Biological Chemistry School of Physical and Mathematical Sciences Nanyang Technological University Singapore, Singapore Anton Wiebe Department of Organic Chemistry Johannes Gutenberg University Mainz, Germany Peter B. Wyatt School of Biological and Chemical Sciences Queen Mary University of London London, United Kingdom Jun-ichi Yoshida Department of Synthetic Chemistry and Biological Chemistry Graduate School of Engineering Kyoto University Kyoto, Japan

Section I Fundamentals and Methods

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1

Basic Concepts Christian Amatore

CONTENTS I. II.

III.

Introduction .............................................................................................................................. 4 Fundamental Aspects of Electron Transfer Reactions ............................................................. 6 A. Oxidation–Reduction Reactions versus Electron Transfer Reactions in Organic Chemistry and Electrochemistry ......................................................................................6 1. Oxidation–Reduction versus Electron Transfer Reactions ........................................ 6 2. Outer Sphere and Inner Sphere Electron Transfers ................................................... 8 B. Standard Potentials: What They Mean and What They Do Not Mean .......................... 11 1. Standard Potentials .................................................................................................. 11 2. Reference Electrodes and Liquid Junction Potentials .............................................. 12 3. Standard Reduction Potentials versus Ionization Potentials or Electron Affinities ................................................................................................................14 4. Comparison or Extrapolation between Solvents ...................................................... 16 5. Formal Reduction Potentials.................................................................................... 17 6. Relationships between Thermodynamic Driving Force ΔE 0 and Feasibility of an Electron Transfer Reaction .............................................................................20 C. Mechanism and Theory of Outer Sphere Electron Transfer Reactions .......................... 22 1. Diffusion Limitation in Electron Transfer Reactions .............................................. 22 2. Basic Features of Electron Transfer in Solution ......................................................26 3. Role of the Solvent: The Outer Shell Reorganization Energy ................................. 31 4. Physical Meaning of the Reorganization Energy and of Energy Diagrams ............ 33 5. Energies and Free Energies in Electron Transfer Theories ..................................... 35 6. Quadratic Free Energy Relationships ...................................................................... 36 7. Cross-Relationships: Evaluation of Rate Constants from Isotopic Rate Constants .....38 8. Experimental Illustrations ....................................................................................... 39 Fundamental Aspects of Electrode Phenomena ..................................................................... 42 A. Monitoring a Half-Reaction: The Electrochemical Cell................................................. 42 1. Monitoring a Half-Reaction ..................................................................................... 42 2. Electrochemical Cell................................................................................................44 3. Resistive Effects and Ohmic Drop: The Supporting Electrolyte–Solvent System ...................................................................................................................... 45 4. Electrode/Solution Interfacial Region......................................................................46 B. General Overview of an Electrode Reaction .................................................................. 47 C. Kinetics of Heterogeneous Electron Transfers ............................................................... 48 1. Mechanism and Theory of Outer Sphere Electron Transfers at Electrodes ............ 48 2. Transfer Coefficient α .............................................................................................. 52 3. Reversible and Irreversible Electron Transfers: Their Role in the Meaning of Oxidation or Reduction Potentials ....................................................................... 53 D. Adsorption Phenomena ................................................................................................... 55

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Organic Electrochemistry

E.

Coupled Chemical Reactions .......................................................................................... 55 1. Electrochemical versus Homogeneous Chemical Reactivities ................................ 55 2. Chemical Reversibility and Irreversibility: Their Role in the Meaning of Oxidation or Reduction Potentials ....................................................................... 57 3. Classification of Coupled Chemical Reactions ........................................................ 59 4. Electrochemical Reaction Mechanisms and the Principle of Microscopic Reversibility .............................................................................................................64 IV. Mass Transfer in Electrochemistry ......................................................................................... 65 A. Fundamental Aspects of Mass Transfer Processes ......................................................... 65 1. Physical Processes of Mass Transfer: Fick’s First Law ...........................................66 2. Relationships between the Electrochemical Current and Mass Transfers............... 68 3. Microscopic Origin of the Diffusion Coefficient: Mass Transfer Rates and Diffusion Layer Thickness....................................................................................... 70 4. Electrochemical Homogeneous Kinetics: Fick’s Second Law................................. 72 5. Dimensionless Formulation of Electrochemical Equations..................................... 74 B. Steady-State Electrochemical Methods: Half-Wave Potential E1/2 ................................. 78 1. Pure Electron Transfer Mechanisms........................................................................ 78 2. Mechanism Involving a Follow-Up Reaction: EC Mechanisms .............................. 81 3. Chemical Kinetics from Half-Wave Potentials: Determination of Rate Constants and Reaction Orders................................................................................84 C. Transient Electrochemical Methods ............................................................................... 87 1. Introduction: Time Hysteresis in Current Reversal Techniques .............................. 87 2. Transient Electrochemical Methods and Chemical Kinetics...................................90 Acknowledgments............................................................................................................................92 References ........................................................................................................................................92 A te convien tenere altro viaggio … Chè questa bestia, per la qual tu gride, Non lascia altrui pasar per la sua via, Ma tanto lo impedisce, che l’uccide… Dante Divina Commedia

I. INTRODUCTION Our objective in this chapter is to familiarize organic chemists with fundamental electrochemical concepts that support several aspects of organic electrochemistry elaborated in this book, as well as many important features of electron transfer reactions. In fact, although the term electrochemistry evokes for most readers the idea of chemistry at electrodes, in our opinion, electron transfer chemistry constitutes a better definition when organic or organometallic electrochemistry is concerned. This is particularly obvious when one thinks of an electrode as a macrosized molecule whose ability to provide or accept electrons is virtually infinite (versus a one-shot molecular redox species) and may be precisely adjusted by fine-tuning of the electrode potential. Similarly, electrochemical reactions involving specific interactions with the electrode material are no more than the analogs of inner sphere, ion-pairing, or complexation reactions. Thus, in many respects, electrochemical reactions do not differ basically from their homogeneous counterparts except for the topology: an electrode is a 2D structure placed in a 3D volume, while molecular redox species are generally dispersed in the solution. This similarity is even more true insofar as it concerns the chemical reactivity of electrogenerated intermediates, since they evolve under conditions that a priori do not differ from those considered in homogeneous chemistry. Indeed, the extraordinarily large electric fields (of magnitude comparable to those at the origin of storm lightning) drop to negligible values within

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5

Basic Concepts

a few angstroms of the electrode surface, in contradiction to generally held ideas. Thus, as soon as molecules have moved over a few molecular diameters from the electrode surface, they probe no special electrical effects associated with their electrochemical origin. This is an important point, although not well recognized or publicized, since it permits an easy transposition of electrochemical results to homogeneous situations or vice versa. Yet there are specific particularities of chemistry at electrodes. They arise from the fact that electrodes both supply and accept electrons to and from molecules dissolved in a solution. This has two important consequences dealing with mass transfer (i.e., transfer of the reagents from the homogeneous volumic space and to the heterogeneous 2D surface) and with current transport across the solution. Transport of the current through the solution requires that rather conductive media be used, at least when high currents are considered. The ensuing necessity for an inert electrolyte is an obvious disadvantage of electrochemical methods, although in our opinion, it is fully compensated by the extraordinary advantage of a precise adjustment of the driving force via the electrode potential. Mass transfer from a volumic homogeneous region, the bulk solution, to a 2D surface results in a spatial structuring of solutions in the close vicinity of the electrode surface. Although such effects are not specific to electrochemistry,* they may be thought of, at first, as additional difficulties to cope with in electrochemistry. However, these spatial structurings may be used with considerable advantage to control the chemical route followed by a given intermediate and force it into a pathway that would never be followed under homogeneous conditions. Indeed, this is used in most electrochemical reactions, especially in indirect reductions or oxidations, in which the proper choice of the electron transfer mediator or of the electrogenerated reagent precursor is crucial to the success of a particular reaction. From our experience, the largest difficulty (or intellectual activation barrier) that homogeneous chemists encounter in trying to deal with the electrochemical literature is directly related to electrochemical jargon rather than to fundamental concepts. Although electrochemistry was born as a synthetic method (by Kolbe, Haber, and Fichter, among others), most electrochemical textbooks use an analytic approach to these fundamental concepts. As a result, mathematical formulations and jargon invade most of the presentations, which often results in an effect similar to the caveat at the entrance of Dante’s Inferno: Per me si va nella città dolente, Per me si va nell’ eterno dolore Per me si va tra la perduta gente … Lasciate ogni speranza, voi ch’ entrate. Dante Alighieri Divina Commedia

Our point is not to criticize these necessary and worthwhile approaches that give to electrochemistry its ability to interpret and rationalize on a quantitative basis a large number of experimental facts that extend far beyond strictly electrochemical domains. Most of our published work (as well as the later sections of this chapter) advocates our belief in the usefulness of these physicomathematical approaches, yet we regret that they may be perceived by nonanalytic chemists as important obstacles in approaching molecular electrochemistry. For this reason, in this chapter, we take advantage of the fact that most of these precise approaches are extremely well exposed in popular electrochemical textbooks [1] to try to present most of the necessary electrochemical basis in * They intervene as soon as heterogeneous reactants or phase boundaries are involved in the reaction medium, for example, as in phase-transfer catalysis. In fact, the conscious (or even subconscious) use of spatial structuring of solutions is not a privilege of electrochemists: nature uses the same phenomenon in most of its reactions, as in the proton transfer pump crucial to ATP synthesis or in the photosynthetic chain.

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Organic Electrochemistry

words and concepts using as much as possible references (or antireferences) to analogous concepts of homogeneous chemistry. Thus, we hope that this chapter may constitute both a whole and a guide to further and more specialized readings in electrochemistry. Yet there are three exceptions, each dealing with kinetic aspects, in which we had to break our resolution. The first two concern electron transfer theories, and the third relates to electrochemical kinetics in the presence of follow-up chemical reactions. Although these three points, and their pertinent derivations, are in our opinion very important for a fine understanding of electrochemical processes and capabilities, we suggest that the reader not interested in the mathematical body may skip the equations but nevertheless follow the corresponding text. This chapter is divided into three parts. In the first, basic definitions and their consequences for homogeneous chemistry are presented. The second deals with the fundamental aspects of electrode phenomena, whereas the third discusses the problem of mass transfer at electrodes and its consequences for electrochemical kinetics. The particular problems and concepts associated with preparative-scale electrolysis are presented in a separate chapter (Chapter 10).

II. FUNDAMENTAL ASPECTS OF ELECTRON TRANSFER REACTIONS A.

OXIDATION–REDUCTION REACTIONS VERSUS ELECTRON TRANSFER REACTIONS IN ORGANIC CHEMISTRY AND ELECTROCHEMISTRY

1. Oxidation–Reduction versus Electron Transfer Reactions In homogeneous chemistry, pure electron transfer reactions are seldom encountered. Indeed, with the exception of a few examples, electron transfers are often associated with atom or group transfers. This usually results in a confused notion of the nature of oxidation–reduction reactions. For example, the reaction of a ketone with sodium in alcohol to afford the corresponding alcohol, via the sequence in the following [2] (see also Chapters 13 and 14), Na C

– ROH O

C

O

Na

C

O–

CH

OH

ROH CH

OH

(1.1)

is termed a reduction. The same term is used for the Meerwein–Ponndorf–Verley reaction, which is supposed to involve a concerted cyclic transition state [3,4]:

AI

AI O

O

O

O

OH

C

C

C

C

CH3

H

(1.2)

CH3

CH3

C

H

CH3

H

Similarly, reactions of metallic hydrides of aluminum or boron, in which the hydride is transferred to the carbon (Equation 1.3 [5], or 1.4 [6,7]), are also considered reductions [7]): Solv C O

Solv

H –Solv

+

+

Li Solv

, AIH–4 Solv

C O

H AIH–2

C

H

O

Li+ Solv

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Solv

H AIH–2

C

H

OH

Li+ Solv

Solv

Solv

Solv

(1.3)

7

Basic Concepts O C O + BH4–

C

ROH

H

O H

C

O

H

– B

R

R H

H

H

(1.4)

O – B

H

H

H

H

Although the preceding four reactions all obey the same stoichiometry for the carbonyl–alcohol transformation and thus involve the same variation in the oxidation state of the carbon atom, they are obviously different. Sodium reduction, in Equation 1.1, supposedly involves single electron [8] and proton transfers in a succession of separate steps; conversely, in reactions (1.2) through (1.4), groups or atoms are transferred. The same confused notions also exist for oxidation reactions, the problem being even more subtle. For example, permanganate is presented in most introductory textbooks as a typical oxidant corresponding to the half-reaction in Equation 1.5: (1.5)

MnO 4 − + 5e + 8H + ⇌ Mn 2+ + 4H 2O

This, associated with the classic one-electron oxidation of ferrous to ferric salts by permanganate, for example, may lead to the implicit notion that permanganate actually sequentially accepts electrons in a similar way as the carbonyl group in Equation 1.1. The fact that this is not always the case in practice is clearly evidenced by the mechanism of alcohol or aldehyde oxidation by permanganate to ketones or carboxylic acids. Indeed, it is generally accepted that the mechanism of acidic oxidation of aldehydes by permanganate involves no electron transfer steps but rather atom transfers as in the sequence of Equations 1.6 through 1.8 (where B− is one of the bases present in the reaction medium) to afford the carboxylic acid and an unstable manganese (V) moiety, which rapidly evolves into MnO2 or Mn2+ according to the pH of the solution [9]. R

R O + H3O+

C

(1.6)

O+ – H + H2O

C

H

H H

H R

C

– O

O+

MnO3

R

C

O

O

MnO3

(1.7) H

H H O

C B–

BH + R – CO2H + MnO–3

R HO

(1.8)

MnO3

Similarly, a reaction proceeding through a sequence of one-electron transfers and chemical steps, such as pinacol formation from a ketone (Equations 1.9 and 1.10), R

R C

R΄

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O

Na/Ng

C R΄

O–

+BH

R C

+B–

R΄

OH

(1.9)

8

Organic Electrochemistry

R

R R

R C

OH +

R΄

O–

C

R

R

C

C

OH

OH

R΄

R΄ C

BH

C

R΄

R΄

(1.10)

R΄ O–

O H

can be reversed via a totally different sequence of steps involving atom transfers as in the classic glycol oxidation to ketones by lead tetraacetate as follows [10]:

OH

C

C

O

C

OH

Pb(OAc)3

+ Pb(OAc)4 C

OH

C

OH

C

OH

C

Pb(OAc)3

C

O

C

O

Slow

O

C

O

C

O

O

(1.11)

Pb(OAc)2 + AcOH

(1.12)

+ Pb(OAc)2

Pb(OAc)2 C

+ AcOH

(1.13)

2. Outer Sphere and Inner Sphere Electron Transfers The preceding examples, which were purposely restricted to well-known reactions of carbonyl and related functions, illustrate the large ambiguities associated with oxidation–reduction notions in organic chemistry. As suggested earlier, these ambiguities certainly stem from the fact that more attention is given to the overall transformation of one of the reactants (the substrate of interest) than to the other(s) (the reagents). Indeed, in the ketone-to-alcohol reductions described, more interest is devoted to the ketone than to the coreactant [Na, Al(OR)3, Li + AlH4−, or BH4−]. When the complete stoichiometrics are considered, it appears more clearly that all four reactions (Equations 1.1 through 1.4) are totally different though substrate and product of interest are formally the same. On the other hand, when one thinks in terms of electrochemical reductions or oxidations, special attention is devoted to the coreactant, that is, to the electrode that provides or accepts electrons. Thus, in order to discuss or compare electrochemical reactions with their organic analogs, it is of the utmost importance to use more precise terms than the so inaccurate reduction or oxidation notions. A similar problem has been addressed in the inorganic and organometallic fields. Indeed, it was early recognized that oxidation–reduction reactions at metal centers must be classified according to two types: outer sphere and inner sphere reactions.* A typical example of the inner/outer sphere dichotomy is given in Equations 1.14 and 1.15, which relate to chromium(II) oxidations by cobalt(III) complexes. * Another complementary concept, named concerted electron transfer, which bridges the gap between homogeneous and electrochemical electron transfers was introduced by Savéant [2d] (see also Chapter 13).

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9

Basic Concepts

Outer sphere (electron transfer) reaction [11]: Cr 2+ + Co( III ) (NH 3 )63+ → Cr 3+ + Co( II ) (NH 3 )6 2+

(1.14)

Inner sphere (ligand transfer) reaction [11,12]: Cr 2+ + Co( III ) (NH 3 )5 Cl 2+ → Cr ( III )Cl 2+ + Co( II ) (NH 3 )52+

(1.15)

Besides their relatively close stoichiometrics, the fact that the fundamental nature of the reactions differs is evidenced by the acceleration by a factor of approximately 6 × 109 in the rate constant when a chlorine ligand is involved in the transition state of the process. Outer sphere electron transfers correspond to situations in which an electron is transferred without the necessity of bond formation or bond cleavage between or within the reactants. The electron is then transferred when the partners are sufficiently close to allow orbitals of suitable geometries to overlap. The energetic stabilization required is of the order of 1 kcal/mol, which is considerably smaller than that corresponding to any bond formation. The ketone reduction by sodium metal in Equation 1.1 is considered to belong to this class of electron transfer. Similarly, owing to the usual poor ability of electrodes in establishing bonds with the electroactive molecules, most of the electrochemical electron transfers pertain to this group.* Interestingly, theories have been developed for outer sphere electron transfers that allow reasonable predictions of their rate constants and activation energies (see Sections II.C and III.C). Note that this definition does not imply that the products obtained upon electron transfer must be stable, but only that the electron transfer activation process does not imply any important molecular rearrangement. Indeed, one or both of the products may chemically evolve through fast follow-up reactions as in the reaction sequences (1.1) or (1.9) and (1.10). Under some specific circumstances, the electron transfer step may be strongly coupled with the chemical reaction expected to follow would a pure outer sphere mechanism be taking place. This is termed concerted (outer sphere) electron transfer [2d]. Inner sphere reduction or oxidation is normally restricted to bond formation between the reactants in the transition state. For example, the chromium (II) oxidation in Equation 1.15 has been shown to involve a chloride bridge between the chromium and cobalt centers as depicted in the following activated complex structure [12]: Cr 2+ ⋯ Cl ⋯ Co( III ) (NH3 )5   

#

(1.16)

The organic chemical reduction–oxidations presented in the reaction sequences (1.2), (1.3), (1.4), (1.7), or (1.11) through (1.13) belong to this class. Obviously, because an electrode is considered able only to supply or accept electrons, this class of reduction–oxidation reactions should not be observed in electrochemistry. However, this is not a clear-cut problem. Indeed, crucial interactions with the electrode surface (e.g., organometallic partial bonding, chemisorption, physisorption) may be crucially involved in the overall electron transfer process although the electrochemical kinetic signature remains characteristic of a classical outer sphere case [13]. In our opinion, a third class of reduction–oxidation reactions exists, which is not clearly encompassed by one of the two just mentioned. It corresponds to the cases in which there really is an electron transfer (i.e., not a group or atom transfer), but the latter is concerted with bond breaking or formation within one of the reactants. It was proposed that the phrase concerted electron transfer * This is not completely true, and a weak chemical interaction with the substrate undergoing the electron transfer step may be crucial. This is always the case in electrocatalysis (e.g., proton reduction, fuel cells) but is hard to observe for organic or organometallic cases. However, there are several perfectly documented examples in the recent literature [13b and references therein].

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Organic Electrochemistry

be used to refer to such situations [2d, 14–16]. A characteristic example of this class is given by the reduction of alkyl halides, represented as follows: (1.17)

RR′R′′C − X + e → RR′R′′C• + X −

Indeed, it is now considered that carbon–halide bond cleavage is concerted with electron uptake (see References 2d,14,15 as well as Chapter 24), that is, there is no intermediacy of an anion radical, − such as [ RR′R′′C − X],• during the electron transfer [15]. These cases must be contrasted with those, such as the following aryl halide reduction: –

Cl + e

(1.18)

+ Cl–, etc

Cl

in which the electron is taken up (or lost) without drastic molecular rearrangement [14,16], to afford an anion radical (or a cation radical in oxidations), although the latter may be extremely short-lived. Formally, the carbonyl reduction in Equation 1.l should fall in this class of reaction, since it is usually considered that the π-carbonyl bond is broken (1.2) upon electron transfer to yield a ketyl anion radical (Equation 1.20) that is better described as a tautomer of the delocalized one in Equation 1.19 than by a simple mesomeric form due to the ion pairing by the alkaline cation: C C

O–

(1.19 and 1.20)

O+e C

O–

Yet the nature of the bond cleaved, especially when the latter is involved in extended delocalizations as in aromatic ketones, and the fact that the skeleton of the molecule remains intact, may explain why these reductions are generally considered as belonging to the outer sphere class. From Table 1.1, which summarizes this discussion, it is seen that outer sphere and inner sphere electron transfers have their exact analogs in both fields of organic chemistry and electrochemistry. TAbLE 1.1 Chemical and Electrochemical Reductions or Oxidations in Organic Chemistry

Designation Outer sphere electron transfer

bond breaking or Formation in the Transition State

Follow-Up Chemical Reaction

Examples Chemical

Electrochemical

No

No

Anthracene → Anthracene

No

Yes

Concerted (outer sphere) electron transfer

No

No

ArCH3 + Fe(III) ⇌ ArCH3 • + + Fe(II) ArCH3−e ⇌ ArCH3• + ArCH3 • + + base → ArCH2•,… ArCH3 • + + base → ArCH2•,… ? RX + e→R• + X−

Inner sphere electron transfer

Within reactants

Yes

Wurtz reaction

Group or atom transfer

Between reactants

Frequent

Meerwein–Ponndorf reaction

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Na

•

Anthracene + e ⇌ Anthracene •

Ag PhCH 2 Cl + e  →…

(see Reference 13b) Invoked to interpret the effect of additives or electrode surface states (as in the Kolbe reaction)

11

Basic Concepts

The third class, that is, concerted (outer sphere) electron transfer, is not easily recognized in organic chemistry, at least to the best of our knowledge. The fourth class, that is, inner sphere group transfers, deals with the peculiar properties of special reactants in organic chemistry or with those of special electrode materials in electrochemistry. At this point, it should be emphasized that there is a large body of outer sphere or inner sphere electron transfer reactions identified and used in electrochemistry, whereas this class of reactions is less developed in organic chemistry. Conversely, there is an extensive variety of chemicals designed for specific group or atom transfer in reduction–oxidation reactions in organic chemistry, but this is an area that is not very developed in molecular electrochemistry. This certainly originates from the obvious fact that the reductive or oxidative strength of an electrode can be varied over a wide range and adjusted with considerable precision because of the easy control of the electrode potential. This advantage is not matched in homogeneous organic chemistry owing to the discrete number of electron transfer reagents, particularly for reductions. On the contrary, surface modifications and reactivities, especially under the conditions of the large electrical fields encountered in electrochemistry (see Section III.4), are not easy to design and control within the present state of the art, whereas owing to the extensive development of organic [7], inorganic, and organometallic [17] chemistry, a large variety of specific agents is made available. However, the apparent disadvantage of electrochemistry in the latter area exists only when the electrode is considered as the only reactant affecting the sought reaction. Indeed, specific chemicals (catalytic or stoichiometric reactants) that affect a desired transformation on a substrate of interest may easily be electrochemically generated or recycled as discussed in this book [18].

B.

STANDARD POTENTIALS: WHAT THEY MEAN AND WHAT THEY DO NOT MEAN

1. Standard Potentials When considering a reaction between a possible electron donor and a possible electron acceptor, it is important to decide a priori whether it will take place or not. This is usually answered by comparison of the standard (reduction) potentials E0 pertinent to each reactant. Indeed, when one considers the possible reaction in the following, +

−

A + D ⇌ A• + D •

(1.21)

the corresponding equilibrium constant K is related to the difference between the standard 0 0 •− (reduction) potentials E A/A and D• + /D couples, respectively. It follows − and E D+ /D of the A/A from the expression of K (F is the Faraday, R is the perfect gas constant and T the absolute temperature) K =e

0 F ( EA − E0 + / A−

D /D

)/ RT

0 0 that at equilibrium, Equation 1.21 should be displaced toward the right-hand side when EA/A − ≫ ED +/D and to the left-hand side when the converse is true (however, see Section II.B.6). When the difference between the values of the two standard reduction potentials is small, the reactant(s) and product(s) equilibrium concentrations obviously depend not only on the potential difference but also greatly on the initial composition of the solution. − According to its definition, the standard (reduction) potential of the A/A • couple is the standard −• electromotive force of a cell in which an A/A electrode is opposed to a normal hydrogen electrode (NHE) whose potential is assigned to zero by convention. −

A+ e ⇌ A• (E 0 )

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(1.22)

12

Organic Electrochemistry

Thus, a standard reduction potential E 0 is ascribed to the half-reaction in Equation 1.22.* The −• potential of the A/A electrode is then given by the Nernst equation [20]: E = E0 +

RT (A) ln −• F (A )

(1.23a)

where (X) represents the activity of species X. Note that in this discussion, a simple or elementary electron transfer is considered. In actual practice, the definition of E0 may be extended to more complex situations (see Section II.B.5) in which the half-reaction includes a series of pre- or postequilibria and at least one electron transfer step. If the pertinent half-reaction is written α1O1 + α 2O2 + ⋯ + ne ⇌ β1R1 + β2 R 2 + ⋯ the potential of the corresponding electrode is given by the Nernst equation: α1

α2

RT ( O1 ) ( O2 ) … E=E + ln nF ( R1 )β1 ( R 2 )β2 … 0

(1.23b)

In practice, there are several limitations to such measurement. Obviously, it implies that both members of the half-reaction are sufficiently stable for a cell to be constructed. This is a serious difficulty in organic chemistry owing to usual great reactivities of the species formed upon electron transfers. For the most frequent cases, it is then impossible to rely on reversible thermodynamic transformations to determine experimental values of standard reduction potentials. However, these important figures, or at least very precisely approximated values, can be obtained from current intensity potential curves or transient electrochemical methods as is discussed in Section III.E. 2. Reference Electrodes and Liquid Junction Potentials Potentials are not measurable on an absolute scale, yet relative scales are easily constructed from measurements of potential differences. This is a fortiori true for electrode potentials. Thus, as explained earlier, electrode potentials are given, by convention, as their potential difference versus the NHE. Yet there are several difficulties associated with the use of the NHE especially under organic conditions. The most obvious is that, although it may be approximated via extrapolations, a real NHE cannot be constructed according to the stipulations included in its definition. The other limitations directly follow from the fact that approximated NHEs are almost impossible to use under situations of organic and organometallic interest. Thus, many other reference electrodes have been proposed and constructed [21]. Each of these electrodes is used by electrochemists according to historical or sentimental reasons or because a given reference electrode is more adequate for a given experimental situation owing to geometric factors or to limit possible pollution of the solution under study. * Because of thermodynamic and electrochemical conventions, standard potentials are defined in the direction of reduction, independently of the respective chemical stabilities of the molecules involved. Thus, for the oxidation of toluene to its cation radical, E 0 refers to the reduction of the highly unstable cation radical into highly stable toluene. To overcome such a priori chemical nonsense, for example, E 0 is frequently designated as the standard oxidation potential of toluene. However, such a term should not be accepted according to canonical rules because it formally implies that the cell now operates in a driven mode, that is, it is connected to an external power supply [19]. Thus, in this chapter, we prefer to use the denomination standard reduction potentials, rather than the usual term standard potential as a reminder of the E 0 definition, although such an expression is basically a pleonasm.

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Basic Concepts

TAbLE 1.2 Potentials of Some Reference Electrodes Reference Electrode Half-Cellb Hg/HgO, NaOH (1 N) Ag/AgCl, KC1 (s) Ag/AgCl, KC1 (1 M) Hg/Hg2Cl2, NaCl (s) Hg/Hg2Cl2, KC1 (s) Hg/Hg2Cl2, KC1 (1 M) Hg/Hg2Cl2, KCl (0.1 M) Hg/Hg2SO4, K2SO4 (s) Hg/HgO, NaOH (0.1 M) a b c d

0 ERef , V versus NHEa

Name

25°C

SSCE SCE NCE

0.14c 0.20c 0.22c 0.236d 0.241d 0.280d 0.334d 0.64c 0.93d

20°C

0.22c 0.245d 0.281d

0 . For any O/R couple, E 0 ( versus NHE ) = E 0 ( versus Ref ) + ERef In aqueous solutions, (s) indicates saturation. From Reference 21c. From Reference 21a,b.

To be considered a suitable reference electrode, an electrode must have a known and reproducible potential versus the NHE. This potential must be nearly an invariant of the current flowing through the electrode, which implies that the electrode reaction is extremely fast and the electrode reactants are extremely concentrated (see Sections III.C.3 and IV.B.2). It should also have a small temperature coefficient and should be easily constructed in a reproducible way. A large variety of electrodes meeting these requirements to different degrees have been devised; Table 1.2 presents a few of the most frequently used. Inspection of Table 1.2 shows that the solution composition of the electrode half-cell is critical. These solutions may differ considerably (nature of the solvent, nature and concentrations of the electrolytes or ionic members of the half-reaction, and so on) from those used in organic or organometallic electrochemistry. The two solutions, that of the experimental system under study and that in the inner reference electrode compartment, then tend to equilibrate across their interface (the liquid junction) to reach a final state in which both solutions are identical. Reaching such mixing equilibrium would lead to pollution of the organic solution on the one hand (i.e., alteration 0 of the reactions of interest) and drift of ERef on the other hand. In practice, to slow down pollution of both solutions significantly, another solution is usually interposed, which acts as a buffer. This bridge then results in the creation of two liquid junctions. At each of these junctions, the physical process of equilibration is controlled by the diffusion of the various components, that is, depends on their mobilities, which may be extremely different [22]. In the process of equilibration, there is then a trend to break the electroneutrality on both sides of the junction. Indeed, the center of positive charges tends to separate from the center of negative charges because of intrinsic differences in the mobilities of the cations and anions. This results in the creation of an electrostatic potential difference Ej across the junction, which opposes the charge separation. The system reaches then a steady-state equilibrium so that the electrostatic energy exactly matches that resulting from differences in chemical composition across the liquid junction. For identical solvents on both sides of the junction, models have been developed to estimate the resulting junction potentials Ej, which may reach several tens of millivolts in practice [23]. Yet in most organic or organometallic experiments, different solvents are used on both sides of the liquid junctions. Other phenomena

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Organic Electrochemistry

(e.g., solvent mixing, different abilities in wetting the physical interface, usually a sintered glass frit or a glass crack, and interfacial and capillary tensions [24]) are then involved. This makes the corresponding analysis intractable or totally irrelevant to real conditions, especially when one considers the difficult reproducibility of most of these phenomena from one experimental setup to another. In practice, it is advisable to use reference electrodes or at least bridges involving, when possible, a solvent identical to that used in the electrochemical cell. Yet this is not always possible, since, for example, other problems may arise from solubilities or temperature dependence. In any case, when potential determinations are crucial, a way to overcome these difficulties is to calibrate the potential scale by measuring the potential of a reference compound (usually ferrocene since solvation of the two redox forms is not expected to vary importantly; see Section II.B.3 and Chapter 7) in the same solution, that is, introducing it in the investigated solution at the end of the experiment. Under such conditions, and provided that the solution composition remains invariant in the course of the experiment, a floating reference electrode may also be used. A silver wire of large surface area or a mercury pool, for example, constitutes such a floating reference electrode. Indeed, it keeps a constant, but unknown potential difference vis-à-vis the solution. Yet the resulting potential scales obtained from run to run need to be reconciled, via the calibration procedure explained earlier, since the floating reference potential is unknown and prone to vary as a function of the exact composition of the solution. 3. Standard Reduction Potentials versus Ionization Potentials or Electron Affinities When one of the members of a given half-reaction is too unstable for the measurement of the pertinent standard reduction potential, one frequently relies on the values of ionization potentials or affinities since these can be measured for a wide variety of species. However, one should be extremely cautious in the use of such data obtained from gas phase when applying them to reactions in solution [25]. Let us consider a simple half-reaction O + e⇌R, with a standard reduction potential E 0. It must be realized that E0 corresponds to (1) stable solvation states for O and R and (2) stable nuclear configurations for O and R. On the other hand, ionization potentials Ip or affinities Aff correspond to nonsolvated O and R and radiative transitions (see Section II.C), that is, to unstable nuclear configurations R* or O*, respectively, for O (affinities) or R (ionization potentials). The ensuing thermodynamic relationships between the three numbers E0, Ip, and Aff are then derived from thermodynamic considerations based on the cycles in Scheme 1.1.* Thus, it follows that E 0 = − Aff +

0 0 0 (∆GO,solv ) − ∆GR*,rel − ∆GR,solv + Cst F

(1.24)

and E0 = Ip +

0 0 0 (∆GO,solv ) − ∆GO*,rel − ∆GR,solv + Cst F

(1.25)

where E 0, Aff, and Ip are expressed in volts, whereas ΔG 0 is in joules and relative to molar quantities. From Equations 1.24 and 1.25, it is seen that E0 and Ip (or E 0 and Aff ) should correlate linearly, * Note that no explicit attention to NHE or metal phases [26] has been given in the cycles, since these factors introduce a constant term Cst in Equation 1.24 or 1.25, independent of the O/R couple. In Scheme 1.1, rel stands for structural relaxation and met1 and met2 for metal 1 and metal 2 to emphasize that generally different metals are used in the different measurements.

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15

Basic Concepts R*gas

NA Aff Ogas+ emet 2

∆G 0R*, rel Rgas ∆G 0R, solv

–∆G 0O, solv ∆G0rad = FE 0

Osolv + emet 1

Rsol

–∆G 0O, solv

∆G 0R, solv Ogas + emet 2 ∆G 0O*, rel

Rgas O*gas + emet 2

–NA IP

SCHEME 1.1 Thermodynamic relationships between ionization potentials, affinities and standard reduction potentials.

with a slope of unity for an extended series of O/R couples, only when the terms in the additional term are reasonably constant within the series. In practice, this condition is almost impossible to fulfill accurately but is approximated for large delocalized molecules belonging to an identical class [27,28]. Indeed, for such molecules, the gain or loss of an electron introduces only small disturbances in the molecules. Similarly, there are nearly no specific solvation effects, the charge being largely delocalized, hence easily screened. Then, provided the equivalent solvation radii are large or remain close within the series, the solvation differences are small [27]. Figure 1.1 shows that when these conditions are fulfilled, good correlations with slopes close to the unity are observed.

Affinities, eV

2

1

–1.5

–0.5

Reduction potentials, V

FIgURE 1.1 Correlation between affinities and reduction potentials for an extended series of polycyclic aromatic hydrocarbons. (Data from Dewar, M.J.S. et al., J. Am. Chem. Soc., 92, 5555, 1970.)

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Organic Electrochemistry

E 0, V vs NHE

2.5

2.0 0.6 V

8.0

1.5

8.5

9.0

Ionization potential, eV

FIgURE 1.2 Correlation between ionization potentials and standard reduction potentials of the corresponding cation radicals for a series of alkylbenzenes. Two dotted lines, with slopes of unity, are positioned at each end of the series to emphasize the deviations vis-à-vis the data in Figure 1.1. (Data from Howell, J.O. et al., J. Am. Chem. Soc., 106, 3968, 1984.)

Thus, determination of E0 for some of the compounds of the series allows good estimates of E 0 to be obtained from Ip or Aff measurements. Yet it must be emphasized at this point that there should be virtually no special difficulties in the direct measurements of the unknown E 0 values for the O/R couples meeting all the requirements discussed here. Indeed, these conditions restrict the validity of the method to chemically reversible outer sphere electron transfers, that is, to those systems for which direct E0 determination does not represent any serious problem (except may be for solubility). From this statement, it is seen that the real practical interest of the method deals with redox couples that involve one too chemically unstable member for experimental E 0 to be determined easily. An example of such a case is given by the series of alkylbenzenes [29] presented in Figure 1.2. Indeed, owing to the rather small size of the benzene π system compared with the large aromatics in Figure 1.1, as well as to the considerable variations in size with the number and nature of the alkyl substituent(s), it seems obvious that the difference in solvation energies in Equation 1.25 has to vary within the series.

( ∆G

0 O,solv

)

0 0 − ∆GR,solv − ∆GO*,rel ≈ −0.3FE 0

Interestingly, a good linear correlation (i.e., with data dispersion similar to that in Figure 1.1) is nevertheless observed in Figure 1.2, yet with a slope of 0.7, instead of unity as in Figure 1.1. From Equation 1.25, it follows that Ip also correlates with E 0, a reasonable conclusion owing to the homogeneity in the series of substituents. Yet the correlation in Figure 1.2 clearly shows the danger of the blind use of such correlations, that is, of the use of Ip or Aff values instead of E0. Indeed, a difference larger than 0.5 V is introduced when extrapolating E 0 from Ip values relative to either ends of the series while assuming a unity slope as in Figure 1.1. 4. Comparison or Extrapolation between Solvents Owing to the extremely large variety of solvents used in organic or organometallic chemistry, it is unlikely that all the pertinent electrochemical data can be found for a particular system and a particular solvent [30]. Thus, the problem of the transposition of data obtained for one solvent to another frequently arises. As discussed for Ip or Aff, this should not be done without extreme precautions.

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Basic Concepts

Indeed, by subtraction of the two equations, akin to Equation 1.25, pertinent to each solvent, one obtains readily for any O/R redox system placed in two different solvents

0

0

E solv1 − E solv2

( ∆G =

0 O,solv1

)

0 0 0 − ∆GO,solv2 − (∆GR,solv1 − ∆GR,so lv2 )

F

+ C′

where C′ is a constant term independent of the O/R couple. However, the numerator of the fraction in the right-hand side, which may include specific solvation terms for each solvent, may vary considerably from one redox couple to the other. Thus, on an absolute basis, extrapolations between solvents are meaningless; however, they may be used with great caution. Evidently, the same is a fortiori true for the use of HOMO or LUMO energies [31] or Hammett–Taft and related extrathermodynamic relationships [32]. 5. Formal Reduction Potentials It is usually difficult to determine or impose activities, since activation coefficients are nearly always unknown for the redox systems investigated in organic or organometallic chemistry. Thus, formal reduction potentials E 0′ are measured, when feasible, rather than E0. E 0′ corresponds to the electrode potential versus NHE when the concentration ratio CO/CR is made unity, rather than the ratio of the activities, as in the definition of E 0 [33]. Thus, from Equation 1.23, it follows that E = E0 +

RT (O) RT f C ln = E0 + ln O O F (R) F f RC R

that is, E 0′ = E 0 +

RT f ln O F fR

(1.26)

where f J is the activity coefficient of species J. The latter being affected by ionic strength and concentrations, for example, it is predicted that E 0′ values should vary from medium to medium. This is important to remember when E 0′ values, determined under the conditions of low concentration and large ionic strength usually met in electrochemistry, need to be transposed under the conditions of low ionic strength and high concentration frequently used in organic chemistry [34]. Formal potential may also include some constant or selected components of the medium that participate in the overall redox reaction. Thus, pH and concentrations of complexing reagents, for example, are often included in E 0′ [35] when the O and R molecules are not interconnected by a simple electron transfer. To explicate this point, let us discuss a typical example related to aldehyde oxidation. The stoichiometry of the reaction in the following involves an overall transfer of two electrons and one oxygen atom: RCHO + 3H 2O ⇌ RCO2H + 2e + 2H 3O +

(E ) 0 1

(1.27)

Yet the acid may exist under two equivalent forms, the acid per se or its conjugated base. Thus, when the pH is such that the carboxylate anion is the stable form, Equation 1.27 is better rewritten as follows: RCHO + 4H 2O ⇌ RCO2 − + 2e + 3H3O +

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(E ) 0 2

(1.28)

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Organic Electrochemistry

From a preparative point of view, both reactions are nearly identical, since the acid and its conjugated base are two forms of the same chemical entity. From a thermodynamic point of view, however, they differ because of the involvement of the acid dissociation in the following: RCO2H + H 2O ⇌ RCO2 − + H 3O +

(Ka )

(1.29)

which thus corresponds to the definition of a second standard reduction potential E20 in Equation 1.28. Yet E10 and E20 are related because of the necessity of a unique potential for the solution. Indeed, from Equation 1.27, the potential is obtained as in Equation 1.30, whereas it is expressed as in Equation 1.31 from 1.28 (note that activity of the water solvent is taken as unity): E = E10 +

RT (RCO2H)(H 3O + )2 ln 2F (RCHO)

(1.30)

E = E20 +

RT (RCO2 − )(H 3O + )3 ln 2F (RCHO)

(1.31)

Identification of both equations readily gives the relationship between E10 and E20 in Equation 1.32, which shows that provided that Ka is known, knowledge of either one of E10 or E20 is sufficient. E20 = E10 −

RT ln K a 2F

(1.32)

On the other hand, owing to the involvement of proton activities in Equation 1.30 or 1.31, it is expected that the acid–aldehyde reduction potential is extremely dependent on the pH. Thus, it is more convenient to separate the pH-dependent terms from those featuring the degree of conversion. Similarly, it is more realistic to consider the sum of the acid and its conjugated base activities to represent the carboxylic product. Thus, taking that into account, Equation 1.30 is rewritten as follows: E = E10 +

(H 3O + )3 [acid] RT RT ln ln + 2 F H 3O + + K a 2 F [aldehyde]

(1.33)

where the term [acid] represents the sum of the concentrations of its two forms, so that E = E 0′ +

[acid] RT ln 2F [aldehyde]

(1.34)

This allows finally the rewriting of Equation 1.33 under the form in Equation 1.34, which directly gives the conversion ratio as a function of the potential and of the pH-dependent formal potential E 0′ in Equation 1.35, where K a′ = K a ( fRCO2H /fRCO2− ) and E10 ′ = E10 + ( RT / 2 F ) ln( fRCO2H /fRCO2− ) whenever activities significantly differ from concentrations (h stands for (H3O+)): E 0′ = E10 +

RT  h3  ln   2 F  h + K a′ 

(1.35)

It is seen from Equation 1.34 that although the significance of E 0′ is not established on definitions as solid as those of E 0, E 0′ nevertheless is of great practical interest since it is the reference parameter that controls the actual conversion ratio in a given situation. Indeed, at any pH or ionic strength,

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19

Basic Concepts

E > E 0′ requires that the acid is the major component in the system, whereas it is the aldehyde for E < E 0′ . At E = E 0′ , 50% of the aldehyde is oxidized. To decide which potential is needed at a given pH, or vice versa, it is then important to know the variations of E 0′ with pH. These variations are conveniently represented in the form of a potential–pH diagram and are derived from Equation 1.35. Indeed, from the latter, it is easily seen that E 0′ varies linearly with the pH as soon as the latter differs slightly from pK a′ , that is (at room temperature), pH < pK a′ pH = pK ′ pH > pK a′

E 0′ ≈ E10′ − 0.06pH

E 0′ ≈ E10′ − 0.01 − 0.06 pK a′ E 0 ′ ≈ E10 ′ + 0.03pK a′ − 0.09pH

Such a diagram is represented in Figure 1.3. The acid predominance zone corresponds to the upper part of the diagram, that is, to pH and potential conditions so that the point representing the system is above the E 0′ versus pH curve. Conversely, the aldehyde predominates when the point figuring the system is located under this boundary. To illustrate the practical interest of these diagrams, let us refer to the classic test of the silver mirror, which corresponds to aldehyde oxidation by the silver nitrate–ammonia reagent. The standard reduction potential for Ag + /Ag is 0.80 V versus NHE. Comparison with the diagram in Figure 1.3 shows that the silver ion is not able to oxidize aldehydes except under basic conditions. Yet under such conditions, silver oxides are formed and precipitated. Ag + + 2NH 3 ⇌ Ag(NH 3 )+ + NH 3 ⇌ Ag(NH 3 )2 +

(1.36)

Thus, a particular base, ammonia, which gives stable complexes with Ag + in Equation 1.36, is used to bring about the basic pH. However, because of the formation of complexes in Equation  1.36,

Acid

E

–1

Aldehyde

–1.5 10

5 pH

FIgURE 1.3 Potential–pH diagram for the aldehyde–acid transformation (solid line). The dashed line gives that for Ag/Ag+ (0.1 M) when the pH is imposed by NH3. The hatched zone corresponds to the only region where the silver mirror test can be performed.

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Organic Electrochemistry

the free Ag+ concentration decreases, which tends to lower the oxidative strength of the solution 0 + ( E = EAg + + 0.06 ln( Ag )). This implies that the ammonia concentration must be adjusted precisely: sufficient for the pH to be in the required range, but not too high in order to retain a sufficient oxidizing power of the silver ion solution. In practice, the best concentration is that corresponding to the titration point of the silver ion, which explains the great caution used in the preparation of the mirror test reagent, [ Ag(NH 3 )2 + ,NO3− ], a fact that is often puzzling for freshman chemists. 6. Relationships between Thermodynamic Driving Force ΔE0 and Feasibility of an Electron Transfer Reaction According to thermodynamics, the feasibility of an electron transfer reaction depends on the free energy change ΔG associated with it. α1O1 + β2 R 2 ⇌ β1R1 + α 2O2

(1.37)

For the electron transfer in Equation 1.37, which involves the two half-reactions in Equations 1.37a,b, α1O1 + ne ⇌ β1R1

(E ) 0 1

(1.37a)

α 2O2 + ne ⇌ β2 R 2

(E )

(1.37b)

0 2

the initial Gibbs energy driving strength ΔGin. is given from basic thermodynamics by*  (R )β1 (O )α2  ∆Gin. = ∆G 0 + RT ln  1 α1 2 β2   (O1 ) (R 2 )  in.

(1.38)

where ∆G 0 = nF ( E20 − E10 ) the subscript “in.” means that the activities to be used in the bracketed term are those corresponding to the initial state of the system. Provided ΔGin. is negative or positive, a reaction proceeds up to when an equilibrium point is reached where ΔG = 0; therefore, at equilibrium, the activities are such as given in Equation 1.39, where the subscript “eq.” indicates equilibrium: 0 0  (R1 )β1 (O2 )α2  = e( nF /RT )( E1 − E2 )  α1 β2   (O1 ) (R 2 )  eq.

(1.39)

From Equation 1.39, it is deduced that for the forward electron transfer in Equation 1.37 to be realized to some extent, it is necessary that the term on the left-hand side of Equation 1.39 is considerably larger than unity, which is supposed to translate into E10 > E20. Conversely, E10 < E20 is often leading to the conclusion that the reaction in Equation 1.37 is not possible. This conclusion may not be true, however, depending on the respective concentrations of the various species considered in Equation 1.37 due to the logarithmic term in Equation 1.38. This is in particular the case when some of the species involved are unstable so that their steady concentrations are always extremely low. * Starting here, we will note (X) the activity of species X. Concentrations, noted [X], may be substituted in all equations for experimental convenience provided the caveats disclosed earlier are duly remembered.

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Basic Concepts

When all reactants and products are stable, these conclusions are valid. However, when one or both products are not stable, Equation 1.37 is continuously displaced to the right, even when thermodynamics would predict that the equilibrium should lie to the left-hand side. This is a situation akin to that usually considered in equilibrium displacements by physical removal (e.g., distillation, precipitation, and complexation) of one of the products. Thus, one may have E10 < E20 and nevertheless observe an efficient redox reaction between O1 and R2. Characteristic examples of such situations are given elsewhere in this book, particularly in Chapters 15, 43, and 44, since this phenomenon is the basis of redox catalysis. In such a situation, the validity of Equations 1.38 and 1.39 is not altered, provided the equilibrium remains established. Yet whenever (R1) or (O2) approaches zero, because of the follow-up chemical reaction(s), the bracketed term on the righthand side of Equation 1.38 remains considerably smaller than unity, opposing the intuitive conclusion based on E10 < E20 . α1O1 + β2 R 2 ⇌ β1R1 + α 2O2 → ⋯

(1.40)

It is thus seen that a direct relationship exists between the thermodynamic driving force ΔE0 and the feasibility of an electron transfer reaction only for those special cases in which all the products and reactants are stable within the time scale considered. Yet it must be pointed out that this is seldom encountered in usual practice owing to the generally high reactivities of the species formed upon electron transfers. This discussion may lead to the extreme conclusion that when the products of an electron transfer reaction are unstable, as in Equation 1.40, thermodynamic figures, such as ΔE 0, are irrelevant. From a simple point of view, this is true. Yet in practice, kinetic notions are implicitly involved in the discussion of the feasibility of any chemical reaction. Indeed, a possible reaction whose completion would require infinite time is not considered feasible. Amazingly, it is because of kinetics that thermodynamic figures recover an interest for situations like that featured in Equation 1.40. To state this point exactly, let us present and discuss the case of the oxidation of methylbenzenes by tris-(l,10-phenanthroline)–iron(III) complexes (FeL33+) in the presence of pyridine bases [36]. From their standard reduction potentials, in the range of 1.35 V versus NHE [37], FeL33+ complexes should not be able to oxidize methylbenzenes, the standard potentials of which exceed 1.75 V versus NHE [29]. The reaction, however, proceeds smoothly in the presence of a pyridine base, according to the following stoichiometry: ArCH3 + 2FeL 33+ + 2py → ArCH 2 − py + + pyH + + 2FeL 32+ and was shown [36,38,39] to initiate through an electron transfer between the iron and the arene centers, as outlined in Scheme 1.2. Two limiting kinetic situations are observed according to

ArCH3 + Fe(III)

ArCH3+ + Fe(II)

ArCH3+ + py

ArCH2 + pyH+z

ArCH2 + Fe(III)

ArCH2+ + Fe(II)

ArCH2+ + py

ArCH2 – py+

(kf, zb)

(1.41)

(kH)

(1.42)

(Fast)

(1.43)

(Fast)

(1.43)

SCHEME 1.2 Mechanistic outline of the oxidation of methylbenzenes by iron(III) phenanthroline complexes.

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Organic Electrochemistry

the respective rates of the proton transfer (Equation 1.42) or of the backward electron transfer in Equation 1.41. When the latter is larger, the electron transfer in Equation 1.41 acts as a rapid equilibrium, the proton transfer being the rate-determining step (RDS). Thus, an apparent rate constant is determined, as follows: k ap = kH

kf kb

(1.45)

where kf and k b are the forward and backward rate constants of the reaction in Equation 1.41. Conversely, when the proton transfer rate is considerably faster than the backward electron transfer, the RDS is the forward electron transfer, and the observed rate constant is then kf, as given here: k ap = kf

(1.46)

It is thus seen that, in the first situation, although the thermodynamic interdiction has been overruled, thermodynamic figures control the apparent rate constant observed. Indeed, Equation 1.45 may be rewritten as follows: k ap = kH e

0 0 ( F /RT )( EFe − EArCH ) 3

(1.47)

Thus, owing to the large magnitude of k H (103–106 M–1 s–1 [36b,c,38,40]), the apparent rate constant in Equation 1.47 remains appreciable even for largely endergonic electron transfers (for 1,2,4,50 0 − EArCH tetramethylbenzene, EFe < −0.6 V). Yet if the electron transfer is too endergonic, kap becomes 3 too small for the reaction to proceed at a significant rate; hence, there is still a thermodynamic limit for the sequence in Equations 1.41 and 1.42, but this limit may be considerably less stringent than a strict application of thermodynamics to Equation 1.41 would predict. The involvement of thermodynamic figures when the forward electron transfer is the RDS is more subtle and arises because of linear or quadratic [41] relationships between activation free energies and thermodynamic driving force, which applies for outer sphere electron transfers like that in Equation 1.41 (see Section II.C.6).

C.

MECHANISM AND THEORY OF OUTER SPHERE ELECTRON TRANSFER REACTIONS

Owing to the aim and scope of the present chapter, it is impossible to discuss here all the subtleties and refinements of electron transfer theories. Yet our purpose is to give a general overview of the different concepts on which these theories are elaborated. As such, this text may constitute per se a presentation of the basic features essential to a thoughtful experimental use of outer sphere electron transfer mechanisms. On the other hand, we hope it may constitute an introduction to more specialized readings [42]. In the following, we tried to rely, for the sake of simplicity, on simple physicochemical pictures that often were not involved in the original works [41,43,44]. As a consequence, this text shortcircuits some extremely important but more specialized theoretical aspects of the original theories. 1. Diffusion Limitation in Electron Transfer Reactions A very common characteristic of electron transfer reactions is that they frequently involve activation energies lower than other chemical reactions that relate to group or atom transfers. As a consequence, the activated processes may be extremely fast compared with the physical process of bringing the two reactants close together (or separating the products) so that they can react. In such cases, the overall rate constant determined does not reflect activation parameters but rather transport properties. Note that occurrence of this phenomenon is also implicitly recognized in homogeneous

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23

Basic Concepts kf

A–+ D+

A+D

Overall (measured)

kb k #f Diffusion (physical)

k#b

r

kpdif

k dif

k act f

Activation (chemical)

[A, D]

[A–, D+]

kbact

SCHEME 1.3 Debye-Smoluchowski microscopic representation of a homogeneous electron transfer indicating its diffusional and activation-driven components.

chemistry. Indeed, the diffusion limit rate constant, of the order of 109–1010 M–1 s–1, corresponds to reactions that are limited by encounters of the reactants, whereas the limit related to product separation is intuitively included within the fuzzy notion invoked in cage separation. Thus, in a general situation, an electron transfer between a donor D and an acceptor A must be considered in terms of, at least, the three successive elementary steps outlined in Scheme 1.3. This representation may be even more segmented when, for example, different ion pairs are involved, such + + − − as [A • , solvent, D •] or [A • , (solvent)n, D •]. Yet it is sufficient here to consider the simplest situation represented in Scheme 1.3 [45]. The formation or dissociation of cages in Scheme 1.3, each a physical process, is normally handled on physicochemical grounds. Yet following a notation by Debye, their r effects may be represented under the form of pseudo-chemical rate constants: kdif (reactant pair) and p kdif (product pair).* Using this notation, it follows [46] that the overall forward kf and backward k b rate constants are given by Equations 1.48 and 1.49 as a function of the activation rate constants kfact and kbact relative to the reacting pairs: 0

f

1 e ∆Gr /RT 1 e ∆G /RT = + + r p kf kfact kdif kdif f

(1.48)

0

1 e ∆Gp /RT 1 e − ∆G /RT = + p + act r kb kb kdif kdif

(1.49)

f f 0 In Equations 1.48 and 1.49, ∆G 0 = F ( ED0 + /D − EA/A − ), while ∆Gr or ∆Gp represents the Gibbs energies necessary to form the pairs in Scheme 1.3 for the reactants and products, respectively. The so-called diffusion rate constants are given by Equations 1.50 and 1.51 [45], where r is the distance between the pair of reacting centers and Dr or Dp the sum of the diffusion coefficients of the reactants or products, respectively.

r kdif =

4πN A Dr

∫

∞

W x / RT x −2e r ( ) dx

(1.50)

r

* Note that this abusive usage (even if correct in terms of physicochemical formulations of rates, but not of rate constants, as recognized by Debye [45]) leads to the fuzzy notions of diffusion and cage separation rate constants.

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Organic Electrochemistry

p kdif =

4πN A Dp

∫

∞

(1.51)

W x / RT x −2e p ( ) dx

r

Both pseudo-rate constants may then be evaluated from a suitable description of the work terms Wr(x) and Wp(x) necessary to bring the reactant centers or those of the products, respectively, from infinity to a distance x. When the work terms are negligible, straightforward integration of Equations 1.50 and 1.51 yields r p kdif = 4πN A Drr and kdif = 4πN A Dpr

which corresponds to the usual diffusion limit rate constants of the order of 109–10l0 M–1 s–1 first proposed by Smoluchowski [45b]. When the work terms are not negligible compared to RT (due to, e.g., specific steric interactions, electrostatic interactions between charged ions, or any other specific ion pairs interactions), the diffusion limit rate constant may differ appreciably from these values [36], being larger when the work terms are negative (attractive pairs) or considerably smaller when the work terms are positive (repulsive pairs). Because of the symmetry in Equations 1.48 and 1.49, we restrict the following discussion to the forward process. It is also noted that kf and k b fulfill the thermodynamic condition in Equation 1.52: kf = kbe ∆G

0

/RT

(1.52)

0′

(1.53)

kfact = kbact e − ∆G

/RT

Indeed, kfact and kbact, being true chemical rate constants, must obey Equation 1.53, where ΔG 0′ is the Gibbs free energy change between the reacting pairs. When there is no other entropic contribution in the formation of the pairs besides those corresponding to the formation of rigid pairs, one has ∆Grf − ∆Gpf ≈ Wr − Wp

and then ∆G 0′ = ∆G 0 − Wr + Wp

From Equation 1.48, it is seen that the observed rate constant is always smaller than or equal to kfdif , defined by considering only the transport limitations in kf expression, that is, by the following equation*: 0

1 1 e ∆G /RT = r + dif p kf kdif kdif

(1.54a)

r p /kdif ), the first term in Equation 1.54a predominates, which means that in When ∆G 0 ≪ RT ln(kdif r the corresponding range, one obtains kfdif ≈ kdif . This corresponds to the traditional diffusion limit, that is, to the case of extremely exergonic reactions. Conversely, for highly endergonic reactions, p −∆G 0 /RT Equation 1.54a yields kfdif ≈ kdif . In the latter case, it might be surprising that a diffusion e control may be observed, yet this apparent paradox corresponds simply to the fulfillment of the thermodynamic conditions in Equations 1.52 and 1.54b, the highly exergonic backward reaction being diffusion controlled. In fact, this later formulation valid for ΔG 0 > 0 provides a formal support for the fuzzy notion of case separation rate constant.

* Note that, similarly, a transport limit is defined for the backward electron transfer. Based on Scheme 1.3 and Equations 1.48 and 1.49, kbdif is given as in Equation 1.54b: kbdif = kfdif e ∆G

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0

/ RT

(1.54b)

25

Basic Concepts

log k dif f

r k dif f ~ k dif

p 0 k dif f ~ k dif exp(–ΔG /RT)

–ΔG 0

0

FIgURE 1.4 Variations in kfdif , as given in Equation 1.54, with the driving force. The shaded area corresponds to the zone where an activation control is not achievable, the overall kinetics being under diffusion control.

Thus, under any circumstances, the overall observed rate constant kf cannot exceed the limit of kfdif given in Equation 1.54a. This limit is represented in the form of a log kfdif versus ΔG 0 plot in Figure 1.4. It must be noticed that since no special considerations have been given to the intrinsic nature of the overall reaction up to this point, these results are valid for any kind of bimolecular chemical reaction. Yet for reactions involving group or atom transfers, Wr(r) and Wp(r) may reach high values (precursor complexes) compared with electron transfer reactions [47]. As a result, the values of p r or kdif may be considerably smaller than the usual figures observed in the 109–1010 M–1 s–1 range kdif for electron transfers. Activation control of the reaction is observed when kf ≪ kfdif , the latter being given in Equation 1.54a. Indeed, formation of the precursor or successor complexes is then considerably faster than the chemical activation process, which therefore limits the overall reaction. This ensues from Equations 1.48 and 1.49 that kf ≈ kf# and kb ≈ kb#, f

kf# = kfact e − ∆Gr /RT

f

and kb# = kbact e − ∆Gp /RT

(1.55)

where the rate constants kf# and kb# defined in Equation 1.55 correspond to the usual measurements of activation rate constants. Indeed, they refer to the activated process “seen” through the diffu+ − sional stages, the initial states being the dissociated A, D or A • , D • pairs, respectively (compare Scheme 1.3). Owing to this formulation, Equations 1.48 and 1.49 are usually rewritten under their equivalent forms as follows: 1 1 1 = # + dif kf kf kf

and

1 1 1 = # + dif kb kb kb

0 kf = e ∆G /RT kb

(1.56a)

(1.56b)

where kf#, kb#, kfdif and kbdif are defined in Equations 1.54a and b and 1.55. Note that these rate constants have exact physicochemical meanings, while kf and k b are operational global values. Though, in general, only kf and k b values are of interest to synthetic chemists. Finally, when using such

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26

Organic Electrochemistry

elegant but compact formulations, it is important to recall that kfdif and kbdif are not the usual diffusion limit rate constants defined in Equations 1.50 and 1.51 but may be considerably smaller when considering endergonic reactions (compare Figure 1.4 and Equation 1.54). In the following, we focus the presentation on the activation contributions, that is, on kf# or kfact, to the macroscopic overall rate constant kf. First, we discuss the nature of an electron transfer reaction. Later, models for the evaluation of the activation energy ∆Gf# are presented. 2. basic Features of Electron Transfer in Solution The fundamental role of the surrounding medium in an electron transfer reaction was early recognized by Libby [48]. Indeed, because of the large number of molecular collisions in solution, molecules are thermalized, in contrast to what occurs in gas phase, though most chemical textbooks rely on gas phase reactions to establish kinetic laws, which are then bluntly applied to solution chemistry. Radiative electron transfers (i.e., analogous to photochemical excitation; see Section II.C.4) do not take place in condensed phases. In other words, an electron cannot be transferred between two systems in which its energy is different, as shown schematically in Figure 1.5a. In the following,+ − a system designates a couple of molecules exchanging an electron, that is, A and D or A • and D •

C

e B

C e

B A

D

D

A

qP

qR

qP

qR C' B'

qelectron

qelectron

A'

(a)

D'

(b)

Electron coordinate A'

B'

e C' qReq (c)

D' q#

qPeq

Nuclear coordinates

FIgURE 1.5 Schematic variation in the electron potential energy as a function of the nuclear coordinates (qR and q P for the reactant and product systems, respectively) and of the electron coordinate (qelectron). (a) Forbidden electron transfer; (b) allowed electron transfer; (c) projection of the system trajectory on the electron coordinate–nuclear coordinates plane. qReq and qPeq are the values of the nuclear coordinate at the equilibrium for the reactant or product systems, respectively, and q# that at the transition state. Solid lines, variations of the nuclear coordinates; (wavy arrow) electron tunneling. Note that, for simplicity, only one nuclear coordinate is represented (see Figure 1.16 and the associated text for a more realistic situation).

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27

Basic Concepts

together with the surrounding solvent. Thus, the initial system corresponds to the reactants and the surrounding medium, in any nuclear configuration, which may greatly differ from its equilibrium state. Similarly, the final system is associated with the products. Because of the intrinsic difference in mass between the nuclei and the transferred electron, the Franck–Condon principle applies during the very act of electron transfer. In other words, nuclear motions (10 –12 s) are then frozen vis-à-vis those of the transferring electron (10 –15 s). Thus, for any nuclear configuration of the initial (or final) system, a potential energy, termed the energy of the initial (or final) system, is associated with the electron, which probes the field of the nuclei (and other electrons), as shown + − schematically in Figure 1.5. Each initial (viz., [A, D]) and final (viz., [ A • , D • ] energy well has a minimum corresponding to the more stable configuration, so that energy increases when the initial and final nuclear configurations depart from these equilibrium positions. Eventually, a point can + − be reached so that [A, D] and [ A • , D • ] have the same nuclear configurations and the same energy. + − At this point only the electron may be transferred, thus converting [A, D] into [ A • , D • ] or reciprocally, thus shifting the system from the initial to the final state, which may then relax to its stable nuclear configuration. However, even when this point is reached (q = q# in Figure 1.5b and c) transfer of the electron (i.e., shifting from the initial to the final energy well) experiences an energetic blockage. Indeed, + − − it is formally stable only “in” D within the [A, D]# state or in A • within the [ A • , D • ]# one. In the physical process of being transferred, the electron must then tunnel through a potential barrier corresponding to the virtual states in which the electron would be between A and D. The height of this barrier is inversely related to the degree of overlapping of the A and D orbitals affecting the electron transfer, that is, of the coupling between the initial and final states at q = q#, which is a point of degeneracy. The intensity of this coupling greatly influences Pe, the probability that the electron tunnels during the time interval in which the initial and final systems have identical energies (compare Figure 1.5b). This time interval, a function of nuclear motions, is in the range of picoseconds, nearly independent of the acceptor and the donor, so that Pe depends mainly on the height of the barrier. As soon as the energy stabilization Vi arising from overlapping orbitals is larger than approximately 1 kcal/mol, Pe is close to unity, and the electron transfer is termed adiabatic. In the converse situation, the electron transfer is called nonadiabatic, and Pe is considerably smaller than unity (see Section II.C.4). From this presentation, it is seen that the overall probability P of electron transfer is tentatively approximated by P = Pn·Pe, where Pn is the probability that the nuclei achieve the q = q# configuration so that the electron energy is identical for the initial and final systems (Figure 1.5b). In usual kinetic terms, this means that the rate constant kfact is given by kfact = κ(kfact )ad, where κ is the transmission coefficient (κ = 1 for an adiabatic electron transfer; κ ≪ 1 for nonadiabatic transfers) and (kfact )ad is the rate constant observed for an adiabatic electron transfer. The latter then depends only on nuclear motions that affect the potential energy of the electron in the initial and final states. In usual chemical terms, (kfact )ad is then directly related to the height of the activation barrier, that is, to the energetic separation ∆Efact between the state where the electron may tunnel and that corresponding to the initial system at equilibrium (compare Figure 1.5b):

( k ) = κ kTh e act f

− ∆Efact /RT

(1.57)

From absolute rate theory [49], kfact is then given by Equation 1.57, where ∆Efact is the activation energy barrier defined usually and expressed for molar quantities. From Equation 1.55, it follows that under such conditions, kf# = κ

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kT − ∆Grf /RT − ∆Efact /RT e e h

28

Organic Electrochemistry

Note that the same presentation applies to kb# and thus kb# = κ

kT − ∆Gpf /RT − ∆Ebact /RT e e h

It is usually more convenient to introduce explicitly the internal energy (Wr and Wp, respectively) and the entropic components of the free energies ∆Grf and ∆Gpf . When the formation of pairs involves no specific interaction, their entropy of formation is identical to that of a rigid rotator (ΔSrot) that yields ∆Grf = Wr − T∆Srot

and

∆Gpf = Wp − T∆Srot

and thus  kT ∆Srot /R −Wr /RT − ∆Efact /RT  kf# = κ  e e e   h  and  kT ∆Srot /R −Wp /RT − ∆Ebact /RT  kb# = κ  e e e   h  From classic statistical thermodynamics [50], the product kT ∆Srot /RT e h is shown to be equal to the collision frequency Z [51], which is usually of the order 1011 M–1 s–1*:  8πRT ( mA + mD )  Z = 10 −3 N A   mA mD  

1/ 2

r2

One then finally obtains [43,44] kf# = κZe −Wr /RT e

− ∆Efact /RT

(1.58)

− / kb# = κZe Wp RT e

− ∆Ebact /RT

(1.59)

and

0′

where kf# = kb# e ∆G /RT because ∆G 0′ = ∆G 0 − Wr + Wp ≈ ∆Efact − ∆Ebact . An absolute expression for κ is nearly impossible to derive in a general case owing to its close dependence on the orbital system of a specific A and D couple [52]. Thus, general theories consider only adiabatic electron transfers, that is, those corresponding to κ = 1. In practice, owing to the very small (less than 1 kcal/mol) interaction required, this is not a strong limitation for most electron transfers in organic chemistry [40]. * In this expression, mA and m D are the molar masses of A and D, and r, the distance between the molecule centers in the precursor complexes, is expressed in angstroms, Z being in M−1 s−1. However, the very physical concept of the collision frequency, Z, is strictly related to gas phase events and has no physicochemical meaning in solution chemistry where molecular trajectories are constantly randomized by shocks with solvent molecules. The expression of Z is nevertheless introduced in condensed phase because ΔSrot value is independent of the way leading to the rigid rotator configuration.

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Basic Concepts

The theoretical predictions of kf# then amount to evaluating the magnitude of ∆Efact in Equation 1.58. Various models have been proposed for this determination, and most of them rely on harmonic approximations in modeling the variations of energy of the initial or final systems with nuclear deformation. The basis of the harmonic assumption is that, for outer sphere electron transfers, (1) only small perturbations in bond length or bond angle are involved and (2) a large number of nuclear coordinates participate in the activation process. Owing to these two considerations, it is understood that although the accumulated energies may be important (from several kilocalories per mole to a few tens of kilocalories per mole), each of the nuclear coordinates may remain close to its equilibrium value, where harmonic descriptions of the corresponding energy variations are roughly accurate. Note in this context that since for inner sphere electron transfers (see Section II.A.2 and Table 1.1) bonds are broken or created in the activation process, the harmonic assumption fails. Indeed, in such electron transfers, the activation energy is concentrated within one given bond, which needs more detailed description of the potential energy reaction coordinate surfaces [54].* Similarly, the geometric requirements for κ to be unity are certainly more important than for outer sphere electron transfers, owing to the extreme localization of the orbitals concerned in the electron exchange. Thus, in the following, we limit our presentation to the harmonic model approximation. Let us denote by qi the ith nuclear coordinate (i = 1, N), qiin and qifin, its equilibrium values in the initial and final systems, respectively, and Ω i, the corresponding reduced force constant (i.e., Ωi = 2Ωiin Ωifin /(Ωiin + Ωifin ), where Ωiin and Ωifin are the force constants in the initial and final systems, respectively [43]). With these notations, the energy of the initial and final systems are given in Equations 1.60 and 1.61: 1 E in = Eeqin + ∑ Ωi (qi − qiin )2 2

(1.60)

1 fin + ∑ Ωi (qi − qifin )2 E fin = Eeq 2

(1.61)

fin where Eeqin and Eeq are the respective equilibrium values, such as

(1.62)

fin Eeq − Eeqin = ∆G 0′ = ∆G 0 − Wr + Wp

The energy E#  of the transition state must be such as E #  = Ein = Efin. Then, by subtraction of Equations 1.60 and 1.61 and denoting the value of qi at the transition state by qi#, one obtains 0 = ∆G 0′ −

1 2

∑Ω ( q i

# i

− qiin

) − (q 2

# i

)

2 − qifin  

(1.63)

Equation 1.63 describes a hypersurface of N − 1 dimensions, which corresponds to all the possible transition states (compare Figure 1.6 when N = 2). Yet the transition state through which the system reacts is that of lower energy. This is obtained by derivation of Equation 1.63 versus qi#. From the resulting equation and Equation 1.63, one shows readily by the Lagrange multiplier method that ∆Efact is given in the following equation: ∆Efact = E # − Eeqin =

λ  ∆G 0′  1 +  4 λ 

2

(1.64)

* Note however that despite this intuitive view, Savéant’s treatment of concerted electron transfers [2c,14–16] (see also Chapter 13) shows that even when extremely nonharmonic contributions are involved, the general formulation derived by Marcus retains its structure.

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Organic Electrochemistry

E

#

A B q1 q2

FIgURE 1.6 Schematic representation of the reaction pathway (…), for two nuclear coordinates (N = 2). The dashed line represents the location of all possible transition states corresponding to the intersection of the two energy wells.

where λ=

1 ∑ Ωi qifin − qiin 2

(

)

2

(1.65)

is the reorganization energy of the system. Because λ is a sum, one may separate the terms relative to bond length or bond angle variations in the A and D molecules (i = 1, M) from those relative to the surrounding medium (i = M + 1,N). Thus, λ is usually written as the sum of two contributions: (1.66)

λ = λi + λo with λi =

1 ∑ Ωi qifin − qiin 2

(

)

2

for i = 1, M

(1.67)

for i = M + 1, N

(1.68)

being the inner shell reorganization energy and λo =

1 ∑ Ωi qifin − qiin 2

(

)

2

the outer shell reorganization energy. Note, however, that albeit the homonymy, these terms must not be confused with the inner sphere or outer sphere nature of an electron transfer (see Section II.A.2 and Table 1.1). λi may be evaluated from x-ray and infrared (IR) data or from theoretical calculations. However, for most organic outer sphere electron transfers, this contribution is usually much smaller than λo because electron transfers involve often very delocalized molecules. In our opinion, one of the

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31

Basic Concepts

greatest merits of the Marcus [43] and Levich–Dogonadze [44] theories is that they allow rather correct predictions of λo through simple equations. Thus, for most outer sphere electron transfers, reasonably accurate values of the rate constants can be predicted. 3. Role of the Solvent: The Outer Shell Reorganization Energy Owing to the usually small energetic interactions between a solvent molecule and the solutes, random fluctuations of the medium surrounding the reactants or the products constantly take place.* Figure 1.7 gives a schematic representation of such solvent fluctuations. During these fluctuations, the solvent nuclei have out-of-equilibrium positions, which then induce nonequilibrium values for the electrostatic field probed by the electron to be transferred. Yet the electrons of the solvent instantly adjust to the local field. This then results in equilibrium values for the electronic components of the electrostatic field probed by the electron exchanged. This dichotomy between the nuclei (vibrational–orientational–translational) and the electronic polarizations arises because of large differences in motion frequencies (10 –12 s for atoms and 10 –15 s for electrons). To determine the resulting energetic variations, the solvent is treated as a dielectric continuum surrounding the reactants or the products. Since we need to consider only virtual charges, the reactant pair (or the product one) is represented by a pair of two metallized spheres† of initial charges ZA = 0 and ZD = 0 (i.e., before the exchange of one electron) for the reactants, or ZA − 1 and ZD + 1 for the products (i.e., after the exchange of one electron). Electron transfer is supposed to proceed through a continuous sequence of virtual states in which only a fractional electron change δ has been transferred, corresponding to virtual charges ZA − δ and Z D + δ for the reactant or product pairs. The metallized sphere radii (respectively, aA and aD in each pair) are supposed constant during the electron transfer, and their centers are separated by a constant distance r. Within this model, the energy associated with a nonequilibrium polarization state is determined by considering that [57] (1) for the solvent nuclei, the polarization corresponds to an equilibrium state in which a virtual fractional charge δe has been transferred, whereas (2) the solvent electronic polarization still corresponds to the real charges. In practice, for the initial system, for example, this is equivalent to determining the variation in the overall electrostatic energy, that is, that associated with the static dielectric constant εs of the solvent when the reactant charges vary from (ZA,Z D) to (ZA − δ, Z D + δ), and to correcting it from the electronic contribution, which is associated with εop, 2 the optical dielectric constant of the solvent (εop = nop , nop being the refractive index of the solvent). From the corresponding energetic cycle represented in Scheme 1.4, it is easily deduced that the ε free energy required to produce an appropriate nonequilibrium polarization is Wδnoneq = Wδεs − Wδ op .

e

(a)

e

(b)

e

(c)

e

(d)

FIgURE 1.7 Schematic representation of the solvent fluctuations inducing the electron transfer. (a, d) Equilibrium states for the reactants and products, respectively. (b, c) Out-of-equilibrium states for the reactants and the products at the transition state. * Note that the tightly bound solvent molecules are then considered as belonging to the inner shell subsystem, and their contributions to the activation energy are then included in λi. † Note that more sophisticated models have been proposed to take into account the nonspherical nature of reactant and products, the eventual distribution of charges on different centers within the molecule [55], the eventual coupling between the inner and solvent fluctuation modes, and other factors [56]. Yet the simplest model described here is sufficient for our purposes, since it includes most of the essential features of electron transfer reactions.

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32

Organic Electrochemistry

Nuclei probe: ZA, ZD Electron probe: ZA, ZD

W δnoneq

Nuclei probe: ZA – δ, ZD + δ Electron probe: ZA, ZD

Equilibrium state

Nonequilibrium state

W δεs

–W δεop

Nuclei probe: ZA – δ, ZD + δ Electron probe: ZA – δ, ZD + δ Equilibrium state

SCHEME 1.4 Thermodynamic cycle representing the relationships between “equilibrium” and “nonequilibrium” states as inferred from Rudolph Marcus concept.

This difference is readily evaluated from simple electrostatics, as in the Born solvation model [27]. ε Thus, both Wδεs and Wδ op are obtained, using εs or εop, respectively, by the sum of the three electrostatic contributions: (1) charge of the A sphere from ZA to ZA − δ, (2) charge of the D sphere from Z D to Z D + δ, and (3) variation of the electrostatic interaction between the two metallized spheres A and D at distance r. Thus, one obtains for ZA = Z D = 0*  N (δe)2   1 1 2 ε Wδεs or Wδ op = −  A + −    8πε   aA aD r  where ε = εs or εop accordingly e is the charge of the electron It is thus seen that for a given δ value, the outer shell component of the energy of the out-of-equilibrium state of the initial system is given by Equation 1.69:  N e2   1 1  1 1 2 2 ( E in )δ = Eeqin −  A   − + − δ   8π   εs εop   aA aD r 

( )

(1.69)

A similar expression would be obtained along the same lines for the final system, yet with δ being replaced by 1 − δ. Comparison of Equations 1.60 and 1.69 shows that the outer shell (solvent) contributions are equivalent to a single harmonic oscillator of force constant Ωo, given as follows:  N e2  1 1 Ωo = −  A  −  4π  εs εop

  1 1 2 + −       aA aD r 

(1.70)

where the virtual charge transferred δ represents the elongation. It then results from Equation 1.68 that λo =

1 1 Ωo (δfin − δin )2 = Ωo 2 2

(1.71)

* When ZA and/or Z D are not null, the establishment of λ o is more complex, but overall, the formulations in Equations 1.69 through 1.72 are retained.

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33

Basic Concepts

since δin = 0 and δfin = 1 for one electron transferred. The outer shell reorganization energy λo is thus finally obtained under the formulation [57] in Equation 1.72:  N e2  1 1   1 1 2 −   + −  λo =  A  a a r 8 π ε ε s  D   op  A

(1.72)

The latter equation, involving easily accessible data, allows suitable predictions of λo for most situations of interest. Note that because of the involvement of a squared term in Equation 1.71, the reorganization energy for a simultaneous transfer of n electrons is n2 times larger than that for a single electron transfer since δfin−δin = n for n electrons transferred. This evidences that in most general situations, electrons are transferred individually and not as multiples. 4. Physical Meaning of the Reorganization Energy and of Energy Diagrams From the preceding discussion, it is understood that although the real phenomenon of electron transfer affects a large number of nuclear coordinates involving the redox centers and a comparatively immense number of solvent molecules, the overall effect is akin to what would be observed for a single generalized coordinate q. This is particularly true for the external medium contribution as established earlier. For simplification of the following discussion, let us assume that only one coordinate q is involved, which varies from qin at the initial system equilibrium to qfin at the final system equilibrium. Let us then introduce the normalized coordinate x = (q−qin)/ (qfin−qin). It is seen from Equation 1.60 or Equations 1.61 and 1.65 that Ein and E fin are given as a function of x by in E in = Eeq + λx 2

(1.73)

fin E fin = Eeq + λ(1 − x )2

(1.74)

with x = 0 corresponding to the initial system equilibrium point, whereas x = 1 corresponds to the final system equilibrium point. Note that the following harmonic model can be adapted to encompass a larger variety of nonharmonic situations that are amenable to harmonic ones simply by using a suitable change of space variable. For example, a simple logarithmic change of variable y = −(ln x)/β, that is, x = exp(−βy), transforms any associative/dissociative Morse potential curve, which represents an energy well in the real space y, into a parabola that represents the same well in the transformed space x. Thus, Equations 1.73 and 1.74 apply in the transformed space, although the real energy well is strongly nonharmonic.* For this reason and for the sake of simplicity, we thus proceed assuming fully harmonic energy wells. Figure 1.8a represents the resulting variations in the two energies with x according to Equations 1.73 and 1.74. It is seen from this figure that, if the electron transfer were to occur radiatively from the initial system equilibrium point, as in the gas phase (e.g., in Ip determinations), then the final system curve would be reached at the point x = 0, fin that is, from Equation 1.74, at E = Eeq + λ. The system would then relax to its equilibrium, that is, fin x = 1 and E = Eeq , thus radiating an energy λ during its reorganization. Such comparison is at the origin of the term reorganization energy for λ though in a real solution system, this path does not occur under thermal activation. A second usual interpretation of the reorganization energy is related to the activation process rather than to a comparison between the gas and the condensed phases. The activation barrier for * In fact, this is the very fact why harmonic approximations, viz., Equation 1.77 keeps its general structure even when electron transfer are concerted with a bond-making or a bond-breaking event, that is, when a fully nonharmonic strong contribution is involved (see concerted electron transfers in References 2c,14,16, see also Chapter 13).

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34

Organic Electrochemistry

E

E

Excited

Fundamental λ 2Vi

ΔG # ΔG 0΄

x

x 0

x

#

1

0

(a)

x

#

1

(b)

FIgURE 1.8 Potential energy–normalized reaction coordinate diagrams, without (a) or with (b) interaction between the initial and final subsystems (see text).

the forward reaction is given by equating Ein and Efin in Equations 1.73 and 1.74. Thus, it follows that x#, the normalized coordinate at the transition state, is given by the resolution of fin E # = Eeqin + λ( x # )2 = Eeq + λ(1 − x # )2 fin which easily yields Equation 1.75 when taking into account that ∆G 0′ = Eeq − Eeqin.

x# =

 1  ∆G 0′ + 1  2 λ 

in ∆Efact = E # − Eeq = λ( x # )2 =

(1.75)

λ  ∆G 0′  1 +  4 λ 

2

(1.76)

Introduction of this value into Equation 1.73 affords the expression of the activation energy in Equation 1.76 for the forward reaction.* Thus, it is seen from the latter equation that λ/4 is equal to the intrinsic activation barrier ∆Go#, that is, to the activation energy that would be obtained for ΔG 0′ = 0. Thus, if a series of electron transfers with virtually identical λ is considered, interpolation of the activation barrier versus ΔG 0′ quadratic law at ΔG 0′ = 0 gives direct access to the experimental value of λ. Proponents of this second interpretation then prefer to reformulate Equations 1.64 or 1.76 under the equivalent form  ∆G 0′  ∆Efact = ∆Go#  1 + #   4∆Go 

2

(1.77)

To conclude, we wish to elaborate more deeply about the energetic diagrams like that presented in Figure 1.8a. As noted previously, the A and D orbitals must interact in order that the electron may be transferred. This simply means that an electronic coupling must occur between the initial and final systems. This results in the mixing of the two systems, which amounts to separating the degenerate * Note that Equation 1.76 is identical to Equation 1.64.

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35

Basic Concepts

potential energy curves in Figure 1.8a into a lower (or fundamental) state and an upper (or excited) state, as represented in Figure 1.8b. Because the potential energies of the electron tend to be identical for the initial and final systems when the transition state is approached, the degree of mixing in this region is higher than near the bottom of each potential well. Thus, the maximal effect is reached at the transition state, where a splitting of magnitude 2Vi is observed [58]. In practice, as soon as 2Vi > RT, that is, Vi > 0.25 kcal/mol at room temperature, when the system crosses the transition point, it has not enough kinetic energy to jump from the fundamental state to the excited one. As a consequence, the system remains on the lower curve and then ultimately reaches the final system equilibrium. This corresponds to a probability of 1 for electron transfer each time the system crosses the barrier (provided its momentum is in the right direction), that is, κ = 1. When 2Vi ≈ RT, when reaching the transition point, the system may jump onto the upper curve. If so, when it goes backward, it falls onto the lower curve, with a large probability (because of the direction of its momentum) of advancing toward the initial potential well. In such a case, the electron has not been transferred and then κ < 1. When 2Vi ≪ RT, this phenomenon occurs almost every time the system crosses the transition point, which means that κ ≈ 0, and no electron transfer occurs. From these considerations, it is understood that Equations 1.64, 1.76, and 1.77, derived without considering electronic coupling, give activation energies overestimated by a term Vi and should normally be corrected. However, as noted for κ, Vi is predictable with difficulty from simple considerations and is usually not included in the activation energy. However, the resulting error is often negligible because λ is not known with a sufficient accuracy [40] (compare the above threshold value of 0.25 kcal/mol for Vi). 5. Energies and Free Energies in Electron Transfer Theories In the preceding description of electron transfer theories, energies were used rather than free energies. For a reaction in liquid phase, ΔE is nearly identical with ΔH, so that the results obtained earlier are valid as soon as entropic variations are negligible,* and thus ΔE ≈ ΔH ≈ ΔG, so that the distinction between ΔG and ΔE may be dismissed. Since outer sphere electron transfers involve small bond lengths or angle variations, this is acceptable for inner shell components of activation free energies. Concerning outer shell components, the approximation is considered not drastic and is equivalent [59] to those made for solvation models (Born) or according to the Debye–Hückel theory [la,c]. Thus, within these approximations and that previously discussed dealing with Vi, it is considered that the activation rate constant of outer sphere electron transfers is given by (from Equation 1.58) kf# = κZe

− ∆Gf# /RT

(1.78)

with ∆Gf# = Wr +

λ  ∆G 0′  1 +  4 λ 

2

(1.79)

where ΔG 0′ is the free energy change between the precursor and successor complexes:

(

)

0 ∆G 0′ = F ED0 + /D − EA/A − − Wr + Wp

(1.80)

and λ is given in Equations 1.66, 1.67, and 1.72. * Note that for the transition state, the entropy corresponding to the formation of pairs is included in the models through the free energies of formation of the precursor or successor complexes using a rigid rotator approximation (see Sections II.C.1 and II.C.2).

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Organic Electrochemistry

6. Quadratic Free Energy Relationships When one considers a series of related electron transfers, that is, involving acceptors (and donors) of reasonably close chemical electronic structures and radii, the reorganization energy λ should remain constant for the series. Thus, it is seen from Equation 1.79 that the activation energy varies as a quadratic function of the driving force, in contrast to the usual observation of linear free energy relationships in organic chemistry (Brönsted correlations between log(k) and log(K), where k and K are the rate and equilibrium constants of the reactions). As a result, the Brönsted coefficient β varies linearly with the driving force, with a value of 0.5 at ΔG 0 = 0. Indeed, from Equation 1.78 and a derivation of Equation 1.79, one obtains  ∂ ln kf# β = − RT  0  ∂∆G

 1 ∆G 0′ = + 2λ  2

(1.81)

When |ΔG 0′| < λ, the predictions of such an equation are in agreement with the deductions of the Hammond postulate. Indeed, for exergonic reactions, β tends to zero when ΔG 0′ → −λ, which corresponds to an activationless reaction. Conversely, for endergonic reactions, β tends to unity when ΔG 0′ → λ, which features a barrierless reaction. However, when |ΔG 0′| > λ, β reaches negative values in the exergonic domain and is greater than unity in the endergonic case, according to Equation 1.81. To recall these facts, the region such as |ΔG 0′| > λ is usually called the abnormal or inverted region, the second term referring to the unusual decrease in rate constant with larger exergonicity of the reaction. In actual practice, there is a large debate about the experimental existence of this inverted region. Moreover, more elaborate models have been proposed involving the possible participation of excited vibronic states, which show that for most organic electron transfers, one should observe a normal behavior, that is [60], ∆G # ≈ Wr , or β = 0 for ∆G 0′ < −λ

(1.82)

∆G # ≈ Wr + ∆G 0′ , or β = 1 for ∆G 0′ > λ

(1.83)

although ΔG #, given in Equation 1.79, remains valid for |ΔG 0′| > λ. Since under electrochemical circumstances, when a conductor electrode is used, the electronic levels in the electrode form a continuum, Efrima et al. considerations [60] fully apply. Hence, no inverted region should be observable, a prediction that is in total agreement with experiments to the best of this author’s knowledge. Similarly, other models, such as the Marcus–Levine–Agmon [61] one, have been developed that also predict the nonexistence of the inverted region. The latter model leads to the ΔG # variations with ΔG 0′ and λ in the following: ∆G # = ∆G 0′ +

0′ λ ln(1 + e − ∆G 4 ln 2 /λ ) 4 ln 2

(1.84)

Particularly interesting is that Equation 1.84 respects the limits in Equations 1.82 and 1.83. For intermediate values of the driving force, the calculated activation barrier from Equation 1.84 is very close to that obtained in the classic harmonic model developed in this section (Equation 1.79). Thus, from an experimental point of view, distinction between the two models is rather difficult or impossible [36a], although seemingly different mathematical expressions for ΔG # variations with ΔG 0′ and λ are obtained.

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To conclude, we discuss the experimental use of Equation 1.81. Indeed, in experimental practice, one has access to overall rate constants kf or k b in Scheme 1.3 and Equations 1.56. Thus, the experimental Brönsted coefficient βexp is determined as follows:  ∂ ln kf  βexp = − RT  0   ∂∆G 

(1.85)

The latter becomes identical to that in Equation 1.81 only when kf ≈ kf# , that is, when kf# ≪ kfdif in Equation 1.54a. However, when diffusion of the reactants or products becomes the rate-determining factor, kf ≈ kfdif , as discussed earlier. Owing to the discreteness of homogeneous experimental data, the shift from an activation control to a diffusion control of the reaction is usually difficult to appreciate, particularly in the endergonic region. The perversity of the problem is even larger when one realizes that application of Equation 1.85 to a diffusion-controlled reaction yields βexp = 0 for exergonic situations and βexp = 1 for endergonic, owing to the two components in the overall pseudo-rate constant kfdif in Equation 1.54a. The difficulty is greater as the reorganization energy λ decreases, as evidenced by Figure 1.9. This caveat [36] is conveniently represented in Figure 1.10,

5 10

log kf

10 20 5

30

0

–10

0

10

–5

ΔG 0΄, kcal/mol

FIgURE 1.9 Variations in the observed rate constant kf as a function of the driving force for various values of the reorganization energy (numbers on the curves in kcal/mol). The solid curves correspond to an activation control. The hatched zone corresponds to a diffusion control (compare Figure 1.4).

20 ΔG 0΄, kcal/mol

I 109 5 . 109 1010 10 5 . 10

20

Activation control

–20

∞, 1011 0

II

0

λ, kcal/mol

0

20

λ, kcal/mol

FIgURE 1.10 Relationship between the driving force and the reorganization energy for the reaction to be under activation control or diffusion control (hatched zone). The boundaries between the two zones are given for various values of krdif = kpdif . The boxes correspond to the location of the experimental systems in Figure 1.11 (I) or 1.12 (II) (see Section II.C.1).

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which summarizes the various conclusions reached here in the form of a plot of ΔG 0′ versus λ. Indeed, it is seen from this diagram that according to the location of a given system, the physical meaning of experimental activation energies obtained by strict application of the following equation (deduced from Equation 1.78, assuming κ = 1) may greatly differ [36a]: # ∆Gexp = − RT ln

kf Z

(1.86)

7. Cross-Relationships: Evaluation of Rate Constants from Isotopic Rate Constants An interesting consequence of the additive formulation in Equations 1.66, 1.67 and 1.68 of the reorganization energy λAD for an A/D electron transfer is that it may be broken in three components: λ oAD , the outer shell reorganization, and λ iA and λ iD, which involve the inner shell contributions related specifically to bond length and bond angle variations in the acceptor (and the donor) center: (1.87)

λ AD = λ oAD + λ Ai + λ Di

On the other hand, for the isotopic electron transfer (i.e., J* is the same species as J except that it is an isotope), −

+

−

+

A* + A • ⇌ A*• + A (or D* + D • ⇌ D∗ • + D) the inner shell component for the reorganization energy becomes λ AA = 2λ Ai i since two A centers are involved; similarly, λ DD = 2λ Di for the isotopic electron transfer relative to i the donor. Thus, it follows that Equation 1.87 may be rewritten under the following form: λ AD = λ oAD +

1 AA λ i + λ DD i 2

(

)

(1.88)

Formally, λ oAD cannot be split into two (A and D) components because of the presence of r ≈ aA + aD in Equation 1.72, which acts as a coupling term. Yet a tedious but elementary algebraic transformation allows Equation 1.72 to be rewritten as λ oAD =

(

1 AA λ o + λ oDD 2



) 1 + ((aa 

− aD )2  2 A + aD )  A

(1.89)

In Equation 1.89, λ oAA and λ oDD are the outer shell reorganization energies for the respective isotopic electron transfers, defined similarly as λ oAD in Equation 1.72. Thus, as soon as the equivalent radii for the acceptor and the donor do not differ by a factor larger than 2, one gets within about 10% precision, that is, perfectly acceptable owing to the crudeness of Equation 1.72: λ oAD =

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1 AA λ o + λ oDD 2

(

)

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Basic Concepts

Introduction of this latter result into Equation 1.88 allows one to formulate λAD as the sum of two components relative to the individual isotopic reactions in Equation 1.90: λ AD =

(

) (

)

(

1  AA  = 1 λ AA + λ DD λ o + λ AA + λ oDD + λ DD i i  2 2

)

(1.90)

The operative interest of such a cross-relationship is extreme, since when associated with Equations 1.78 through 1.80, it allows an a priori estimation of the rate constant of any outer sphere electron transfer reaction provided the corresponding standard reduction potentials and isotopic rate constants are known. For kinetics, this is the equivalent of E 0 for thermodynamics, since knowledge of a series of half-cell potentials (E 0) allows the voltage of any battery formed by combination of two such half-cell to be evaluated. Yet a caveat in this approach is that it should be ensured that κ = 1 for all reactions, that is, there is always sufficient overlap between the orbitals [62]. 8. Experimental Illustrations The purpose of this section is not to present a compendium or a discussion of the numerous experimental illustrations of the validity of electron transfer theories, but rather to select two typical examples that illustrate the main features of these theories in organic chemistry [63]. a. Validity of QFER over a Wide Domain of Free Energies The main difficulty in discussing the dependence of rate constants with driving force over a wide domain stems from the fact that, to be significant, the series must involve electron transfers with nearly constant reorganization energies. However, it is easily understood that the two requirements (ΔG 0 variations over a range of several tens of kilocalories per mole but λ invariant) are difficult to satisfy simultaneously. Such a series, which extends over nearly 70 kcal/mol in ΔG 0, has been presented by Rehm and Weller [64]. The series corresponds to electron transfer between an excited acceptor and a donor acting as a quencher or between an excited donor and an acceptor quencher. Thus, rate constants, ranging from 106 to 2 × 1010 M–1 s–1, have been determined for more than 60 different donor–acceptor pairs in acetonitrile. Examination of the corresponding results, presented in Figure 1.11, clearly shows that three domains may be defined: (1) ΔG 0 < −10 kcal/mol, where kf is independent of the driving force, that is, where βexp = 0; (2) −10 kcal/mol < ΔG 0 < 10 kcal/mol, where βexp varies from zero to unity; (3) ΔG 0 > 10 kcal/mol, where β = 1. These data were corrected for diffusion components, so they correspond to an activation control of the electron transfer reactions.

log kf

10

8

0

–10

–20

–30

–40

–50

ΔG 0, kcal/mol

FIgURE 1.11 Linear energy relationship over an extended domain of driving force. (Experimental data from Rehm, D. and Weller, A., Isr. J. Chem., 8, 259, 1970.)

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They qualitatively agree with the predictions of Equations 1.78 and 1.79 for |ΔG 0| < 10 kcal/mol but show that Equation 1.79 is not valid outside this range, Equations 1.82 and 1.83 being more adequate. As such, they have been considered a convincing experimental argument for the nonexistence of the inverted region [65].* Yet, owing to the large scatter of the data, a precise determination of the validity of Equation 1.79 in the intermediate (|ΔG 0| < 10 kcal/mol) free energy range is impossible. b. Validity of the QFER for Low-Endergonicities or Exergonicities In order to test the experimental validity of the quadratic relationship between ΔG # and ΔG 0′, a rather small free energy domain needs to be investigated. As a consequence, even small variations in λ result in too large a scatter of the data for significant quantitative conclusions to be drawn (compare the data in Figure 1.11 for |ΔG 0| < 10 kcal/mol). Thus, a series of very closely related redox couples must be investigated.† A series meeting these requirements is given by the endergonic oxidation in acetonitrile of methylbenzenes by tris-(5X-phenantroline)-iron(III), X = H, Cl, NO2, already discussed [36a,c]. Indeed, for such a series, the inner shell contributions are minimal owing to the delocalization and close identity of the iron(II/III) complexes [66] and the small distortions of the arene cation radical vis-à-vis the neutral (for C6H6 + •, the inner shell reorganization energy from a nuclear configuration corresponding to benzene involves a Jahn–Teller distortion [67] and has been evaluated to be approximately 3 kcal/mol by Salem [67c]). Thus, the outer shell reorganization energy predicted according to Equation 1.72 is (λ o in kcal/mol: aAr, aFe, and r = aAr + aFe in Å)  1 1  1 1 2 λ o = 1.67 × 102  −  + −  ε ε a a r s   Ar Fe  op * However, it should not be concluded that the inverted region does not exist. A visible proof of its existence when the density of rovibronic states is scarce is given by photosynthesis. In plants, the inverted region process impedes the recombination of the electron/hole pair created by photochemical process, so that a much less favorable process (thermodynamically speaking) can occur (e.g., reduction of quinone carriers). In fact, the inverted region simply shows up when a too large amount of energy (generated by an extremely downhill electron transfer) is dissipated for a few picoseconds. When the molecules engaged in this process are too rigid to absorb this large Gibbs energy release within the required time range, the system cannot proceed. This is reflected by an increasing activation barrier when ΔG 0 is made more and more negative beyond ΔG 0 = −λ. However, if ΔG 0 is negative enough that rovibronically excited states of the final system can be met, then the reaction proceeds through such states. Inverted region (fundamental-fundamental) R

Non inverted region: ECL (fundamental-vibronic-fundamental) P

R

P*

P

# #

hv

†

These can in turn relax to the fundamental final system state either by the release of a visible photon (which is the basis of electrochemiluminescence or ECL) or a cascade of infrared photons generating heat. In Rehm and Weller work [64], system pairs possessed large densities of rovibronic excited states so that the inverted region could not be observed [60]. Note that the corresponding condition is more easily satisfied for electron transfer at electrodes, owing to the easy variation of the driving force, by adjustment of the electrode potential without modification of the redox couple.

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Basic Concepts

5

0.25

ΔG #/λ

log kf 0.50 0

ΔG 0, kcal/mol

0.75

ΔG 0/λ

–5 20

10

0.75

0

(a)

0.50

0.25

0

(b)

FIgURE 1.12 Quadratic free energy relationship (QFER) in the low endergonic region for the oxidation of methylbenzenes by Fe(phen)33+ . (a) Experimental variation of the observed rate constant compared with the theoretical variations (solid curve): (- - -) activation or (−⋅−⋅−) diffusion controls, (b) experimental relationship between the activation free energy and the driving force compared with the theoretical variations for λ = 27 kcal/mol in the solid curve: (- - -) polynomial regression in Equation 1.92 [36a,c]. (Data from Schlesener, C.J., J. Am. Chem. Soc., 106, 3567, 1984; Schlesener, C.J. et al., J. Am. Chem. Soc., 106, 7472, 1984; Schlesener, C.J. et al., J. Phys. Chem., 90, 3747, 1986.) 2 = (1.34)2 for that is, λo ≈ 21 kcal/mol for aAr ≈ 3.5 Å, aFe ≈ 7 Å [68], and εs = 37.5, and εop = nop acetonitrile. The resulting predicted value of the global reorganization energy is then predicted to be ca. 24 kcal/mol for each couple in the series. Figure 1.12a indicates how the measured rate constant kf for the endergonic electron transfer oxidation of methylbenzene by (phen)3-iron(III) complexes varies with the driving force. It is seen from these data that, provided that ΔG 0′ < 10 kcal/mol, a very good fit is obtained with the rate constant predicted from Equations 1.78 through 1.80 on the basis of an activation control of the overall reaction when using κ = 1, Z = 1011 M–1 s–1, and λ = 24 kcal/mol, determined earlier. However, for larger positive values of ΔG 0′, the diffusion cage processes contribute significantly to the experimental rate constant. Yet the activation rate constant kf# may be extracted from the experimental kf data through the use of Equation 1.56a. The real activation free energy ∆Gf# is then derived through Equation 1.78 and plotted as a function of ΔG 0′ = ΔG 0 + Wp in Figure 1.12b. Quadratic regression analysis of these data allowed their fit by a second-order polynomial expression in Equation 1.91 [36a,c] (represented as the dashed line in Figure 1.12b):

∆Gf# = 6.7 + 0.50∆G 0′ + 8.7 × 10 −3 (∆G 0′ )2

(1.91)

which is in good agreement with the predicted variations in Equation 1.79 for λ = 27 kcal/mol. Indeed, development of Equation 1.79 with λ = 27 kcal/mol yields ∆Gf# = 6.7 + 0.50∆G 0′ + 9.3 × 10 −3 (∆G 0′ )2

(1.92)

Another way to appreciate the correctness of Equation 1.79 in the prediction of the experimental barrier consists in the evaluation of λ for each individual couple as presented in Table 1.3. Indeed, from Equation 1.79, λ is obtained through the solution of the second-order equation

(

)

λ 2 + 2 ∆G 0′ − 2∆Gf# λ + (∆G 0′ )2 = 0

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(1.93a)

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Organic Electrochemistry

TAbLE 1.3 Experimental Evaluation of the Reorganization Energy λ for the Oxidation of Methylbenzenes by Tris-(5X-phenantroline)–Iron(III) Complexes [36a] Methylbenzene

a b c

X

ΔG0’a

∆Gf#a

λa

λava,b

H Cl NO2

11.2 8.86 6.55

13.4 12.6 10.7

26.4 30.1 28.2

28.2 (1.8)

H Cl NO2 H C1 NO2 H Cl NO2

14.2 11.9 9.55 16.0 13.7 11.4 15.8 13.5 11.2

15.6 13.9 12.4 16.8 13.4 13.3 18.0 14.0 14.0

26.4 26.4 27.1 24.9 —c 25.2 32.8 19.8 29.3

26.3 (0.4)

25.1 (0.2)

27.3 (6.7)

In kcal/mol. Average value for a given arene, standard deviation indicated in parentheses. For all couples, λav = 27.0 ( ± 3.3) kcal/mol. Not determinable since ΔGf > ΔG0′.

∆Gf# = RT ln

Z kf#

(1.93b)

where ΔG 0′ and the experimental activation free energy ∆Gf# from Equation 1.93 are known. Thus, a value of 27.0 kcal/mol (standard deviation 3.3 kcal/mol, 11 data) is obtained for the average reorganization energy, which is in very close agreement with that predicted (24 kcal/mol).

III. FUNDAMENTAL ASPECTS OF ELECTRODE PHENOMENA A.

MONITORING A HALF-REACTION: THE ELECTROCHEMICAL CELL

1. Monitoring a Half-Reaction In homogeneous chemistry, an electron transfer reaction involves two reactants, an acceptor A and a donor D. Thus, formally, an electron transfer reaction consists of the superimposition of two elementary chemical acts: the reduction of the acceptor and the oxidation of the donor. Yet as discussed in the first part of this chapter, since more attention is given to one of these elemental reactions, the overall reaction is termed a reduction when the acceptor is the compound of interest or an oxidation when chemical transformation of the donor is favored. In electrochemistry, the same phenomenon (essentially related to charge conservation) occurs, yet the reduction of the acceptor A occurs at one electrode (the cathode in electrolytic cells) and the oxidation of the donor D at the other (anode). Thus, the kinetics of the overall cell reaction* depends on both half-reactions, in a similar way as the kinetics of a homogeneous electron transfer depends on the acceptor and the donor. However, because the two half-reactions occur at different locations, it is technologically possible to break the coupling between the two electrodes. * That is, the current flowing through the cell, since it corresponds to an identical number of electrons consumed at each electrode per unit of time.

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Basic Concepts W

I

φan

R

Ri

Ccat

φcat (a)

Zcat

(c)

Potentiostat E + R0i

E'

E + R0i Z cat

Can

Ccat

Can

(b)

R

Zan

(d)

Ccat

Ref

A

E 0Ref R0

R – R0

Zn

Can

E Potential generator

Ref W

A

(e)

FIgURE 1.13 Representation of the potential variation in the electrochemical cell and associated electronic schemes (a–d). (Φcell = Φcat + Ri + Φan; see text). W, working electrode; Ref, reference electrode; A, auxiliary electrode. (e) Schematic description of the electrochemical cell with a potentiostat.

This is done via two different approaches, based on the same principle, that take advantage of the potential distribution in an electrochemical cell, shown in Figure 1.13a. From this schematic representation, it is seen that the potential difference Φ between the two electrodes is the sum of three components: two of them, Φcat and Φan, are related to each electrode, and the third, Ri, arises because of the ohmic drop due to the current intensity, i, flowing through the cell solution resistance. The potential variations Φcat and Φan are localized within an extremely thin layer ( 10) and supporting electrolytes with rather large ions (e.g., NBu4 + BF4−). Considering these conditions and that the solvent should be able to dissolve organic molecules explains why most electrochemical experiments involve a narrow class of solvents, such as acetonitrile (εs = 37.5), dimethylformamide (40), dimethyl sulfoxide (46.7), and propylene carbonate (36.7) [71]. 4. Electrode/Solution Interfacial Region To simplify the following discussion, we restrict it to the case of a cathode. Obviously, the results and phenomena directly apply to anodes, cations being replaced by anions, and vice versa. Experimentally, it was early recognized that an electrode/solution interphase behaves as a capacitor† [74]. This is because the negatively charged surface of the cathode generates a very strong electrical field that tends to attract positive ions from the solution, as sketched in Figure 1.14a. The positive layer thus formed exerts, on the solution side, an electrical field of opposite direction that attracts negative ions from the solution. However, since this latter field is partly compensated by that resulting from the negative charges in the electrode, the charge density of the anionic layer is smaller than that of the cationic one. Repetition of this phenomenon results in alternating layers of positive and negative charges, with a charge density in each layer decreasing with the electrode distance, as pictured in Figure 1.14b. Because of the decrease in electrical field with distance, the successive layers are less and less bound. Then, owing to thermal agitation, these layers cannot be as rigid as might be deduced from the preceding presentation and interpenetrate each other except for the first, which is tightly bound * Inert means that the electrolyte is not involved in the electrochemical transformation occurring at the electrodes. The electrolyte is usually also called the supporting electrolyte after its function of supporting the current flow through the cell. † Note that this property is used in electronics in the so-called electrochemical capacitors, which are no more than electrochemical cells without faradaic reactions within their potential range of use. That is why these capacitors have a polarity and a voltage limit; they must then be used with non-alternative voltages to avoid their explosion due to undue Faradaic production of gazes.

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Basic Concepts

φM

OHP

x φsol (a)

(b)

(c)

0 (d)

xOHP

x0

FIgURE 1.14 Schematic representation of the electrode–solution interfacial region. (a) Helmholtz model; (b) structured layer model; (c) thermally disorganized layers; and (d) resulting potential variations with distance from the electrode; ΦM, electrode potential; Φsol, solution potential; OHP, outer Helmholtz plane (few Å); x0, extremity of the diffuse layer (few tens of Å); x < xOHP, compact layer; xOHP < x < x0, diffuse layer.

onto the electrode or onto a tightly bound strong dipolar layer of solvent [75]. This phenomenon is indicated by the corresponding names compact layer for the first layer and diffuse layer for the remaining part; both of them constitute the double layer and act as two capacitors in series. The capacitance (usually 20–50 μF/cm2) and dimensions (a few angstroms for the compact layer and a few tens of angstroms for the diffuse layer) may be evaluated via the Gouy–Chapman–Stern theory [76], which is developed in most electrochemical textbooks. The main conclusions of this model, which are in general agreement with experimental observations, are the following: (1) because of the compact layer, a large potential drop arises within a few angstroms of the electrode surface (compare Figure 1.14d); thus, most of the potential difference between the electrode and the solution is lost at the end of the compact layer (i.e., at the outer Helmholtz plane, OHP); (2) the potential difference between the OHP and the solution results in an electrical field considerably lower than that between the electrode and the OHP, which nearly exponentially decays with distance; (3) beyond several tens of angstroms from the electrode, the potential is identical to that in the bulk of the solution, except for eventual ohmic drop effects, as discussed earlier (compare Figure 1.13a); (4) as a result of the third conclusion, beyond a few tens of angstroms, the electrode exerts no electrostatic attraction or repulsion on any ionic species, which explains why an anion may be reduced or a cation oxidized at electrodes.

B.

GENERAL OVERVIEW OF AN ELECTRODE REACTION

When thinking of any electrochemical reaction, it is important to remember that the very act of electron transfer takes place at a surface, whereas the material to be reduced or oxidized is dispersed in a volumetric phase. Thus, mass transfer from the bulk of the solution (or to the bulk solution, for the products) plays a central role in electrochemical processes. As such, the physical processes of mass transfer (diffusion, migration, and convection) have received considerable attention in the electrochemical literature and are discussed separately in Section IV.A.1. For our purpose here, it is sufficient to know that mass transfer results in the creation of concentration profiles as a function of the distance from the electrode, as illustrated in Figure 1.15. For most of the usual electrochemical situations, these profiles extend over a few to a few tens of micrometers from the electrode surface. Beyond this region, called the diffusion layer, where mass transfer processes are effective, the solution is homogeneous (except for microscopic convective effects as in homogeneous chemistry). Thus, a simple electrode reaction involves successively the following three steps: (1) mass transfer of the reactant from the bulk solution to the electrode, (2) electron transfer at the electrode

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Organic Electrochemistry

CR CR

CP 0 (a)

CP

x δ

δconv

0

x δconv

(b)

FIgURE 1.15 Concentration profiles for the reactant and the product as a function of the electrode distance x in (a) transient or (b) steady-state electrochemical techniques: δ and δconv, diffusion layer and stagnant layer thicknesses.

surface, and (3) mass transfer of the product of the reaction from the electrode surface to the bulk solution. These three elementary steps may be perturbed by earlier or later chemical steps and by adsorption or desorption phenomena at the electrode surface. Each of these processes may become the RDS or may interfere significantly in the overall kinetics. At first glance, the situation then seems much more complicated than in homogeneous chemistry. Yet this is only an illusion because most of these steps have their exact counterparts in homogeneous chemistry. Mass transfer, though occurring here over macroscopic dimensions, corresponds to the diffusion control of a chemical reaction as described. Earlier or later chemical steps are involved in most homogeneous electron transfer reactions. Adsorption and desorption are analogous to the formation of stable precursor or successor complexes. Electron transfer to or from an adsorbed species is very reminiscent of inner sphere reductions or oxidations. These analogies are not fortuitous; indeed, an electrode may be viewed as a gigantic molecule with an infinite supply of reducing or oxidizing power adjustable at will.

C. KINETICS OF HETEROGENEOUS ELECTRON TRANSFERS 1. Mechanism and Theory of Outer Sphere Electron Transfers at Electrodes In the following, we consider an outer sphere electron transfer reaction involving the exchange of n electrons (n > 0 for a reduction; n < 0 for an oxidation), taking place at an electrode whose potential is E*: R + ne ⇌ P (E 0 )

(1.96)

The electrode is supposed to have no particular interaction with R or P. Its surface area is hereafter designated by A. Because in electrochemical reactions electron transfers occur near the electrode/solution interface, that is, at a 2D surface, it is more suitable to refer to the number of moles converted per unit of time and unit of electrode surface area, dNR /dt and dNP/dt, rather than using concentrations as in homogeneous kinetics. Since there is no stable interaction with the electrode, dNR /dt = −dNP/dt. * From the discussion in Section III.A.1, E = ΦW − ΦRef. Yet when E 0 is given vis-à-vis the same reference, the contribution of the reference cancels in the difference E − E 0. Thus, in the following, E is considered the electrode potential vis-à-vis that of the solution.

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Basic Concepts

R0 + ne

Overall:

kf kb

P0

(E0,ΔG0)

PΦ

(E0 + Φs, ΔG0')

Diffusion/migration:

Activation:

RΦ + ne

k act f k act b

SCHEME 1.5 Debye-Smoluchowski microscopic representation of a heterogeneous electron transfer indicating its diffusional and activation-driven components (compare to Scheme 1.3 for a homogenous situation).

On the other hand, each mole of the reactant R electrolyzed corresponds to a charge nF transferred that affords the charge consumed per unit of time, that is, the current: i = −nFA

dN R dN = nFA P dt dt

(1.97)

Since the definition of E 0 in Equation 1.96 refers to standard conditions, that is, conditions in which no external electrical field applies on the reactant or the product, it is more convenient to consider the act of electron transfer as a succession of three individual steps as outlined in Scheme 1.5, which is reminiscent of that in Scheme 1.3 established for the homogeneous analogous situation (see Section II.C.1). In Scheme 1.5, R0 or P0 relates to R or P in the closest plane to the electrode where no electrical potential applies, that is, at the end of the diffuse layer,* denoted x0 in Figure 1.14d. RΦ and PΦ relate to R and P at the site of electron transfer xΦ, usually considered close to or slightly within the OHP, where an electrical potential Φs (i.e., the electrical potential at the electron transfer site) applies. Let (R)0 and (P)0 be the activities of R and P at x0 and (R)Φ and (P)Φ those at x = xΦ (the electron activity is not considered because it is implicitly included in the potentials). For all the electron transfer reactions investigated up to now, the rates of the diffusion–migration steps in Scheme 1.5 do not limit the overall process. Although this may be of importance for the prediction of the maximal rate of electron transfer at an electrode, we do not consider the possible kinetic limitation by either of these physical processes (compare the diffusion control of a homogeneous electron transfer discussed earlier). Thus, within this restriction, the existence of the diffusion–migration processes reflects only thermodynamic contributions. Thus, the overall rate constants in Scheme 1.5, defined in Equation 1.98, i −dN R dN P = = = kf (R )0 − kb (P )0 nFA dt dt

(1.98)

are related to the true activation constants via the relationship in Equations 1.99 and 1.100, where ∆GΦR and ∆GΦP are the respective free energy changes for R and P between the end of the diffuse layer and the electron transfer site. R

kf = kfact e −∆GΦ / RT

(1.99)

* Note that this limit must not be confused with that of the diffusion layer/bulk solution interface, owing to the extreme differences in size: some tens of angstroms for the double layer and some tens of micrometers for the diffusion layer.

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Organic Electrochemistry P

kb = kbact e −∆GΦ / RT

(1.100)

Note that because of their definitions in Equation 1.98, the heterogeneous rate constants are expressed in units of length per units of time (usually in cm/s in the electrochemical literature). As for the homogeneous case, it is more convenient to separate the terms arising from electrical potential from those, ΔSpair, of entropic nature, due to the formation of a pair electrode/molecule. Thus, ∆GΦR = Z R FΦ s − T∆Spair and ∆GΦP = Z P FΦ s − T∆Spair, where Z R and Z P are the respective ionic charges of R and P (note that Z P = Z R − n owing to the charge balance). On the other hand, kfact and kbact are obtained from the absolute rate constant theory as kfact =

kT −( ∆Gfact /RT ) e h

kbact =

kT −( ∆Gbact /RT ) e h

This allows one finally to write the experimentally observed rate constants as follows: kf = Z ele − Z R FΦ /RT e

− ∆Gfact /RT

(1.101)

kb = Z ele − Z P FΦ /RT e

− ∆Gbact /RT

(1.102)

where Z el = (kT/h)e∆S / R is shown from classic statistic thermodynamics to be related to the collision frequency of a molecule on a wall [50], that is, pair

1/ 2

 RT  Z el =    2πm 

(1.103)

in which m is the molar mass of R or P. For most organic molecules, Z el is of the order of approximately 2 × 103 cm/s. Note, however, that no transmission coefficient (compare the κ terms in the homogeneous analog rate constants in Equations 1.58 and 1.59) is included in Equations 1.101 and 1.102. Indeed, the orbital variety at an electrode surface is sufficiently wide for convenient overlap to occur with most organic molecules prone to outer sphere electron transfer, so that the adiabatic requirement is generally fulfilled. Moreover, Hale [77] estimated that electron transfers at an electrode are adiabatic as soon as the electron transfer site is within 15 Å of the electrode surface, for most organic molecules. Thus, determination of the rate constants in Equation 1.101 or 1.102 requires only that of ∆Gfact since ∆Gbact = ∆Gfact − ∆G 0′ , ΔG 0′ being the free energy of the electron transfer between RΦ and PΦ (compare Scheme 1.5):

(

)

∆G 0′ = ∆G 0 − Z R FΦ s + Z P FΦ s = nF E − E 0 − nFΦ s that is,

(

∆G 0′ = nF E − E 0 − Φ s

)

(1.104)

The forward activation free energy is determined along the same lines as presented for homogeneous electron transfer since both problems are strictly equivalent [43,44,78,79]. Thus, ∆Gfact is given

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as a function of the local driving force (ΔG 0′ in Equation 1.104) and of the associated reorganization energy λ el by the expression in Equation 1.105, which derives from Equation 1.64. ∆Gfact =

λ el  ∆G 0′   1 + el  4  λ 

2

(1.105)

Again, λ el involves two contributions, of inner shell λ eli and outer shell λ elo nature, but these contributions relate only to the R/P couple since the electrode is not affected by the electron transfer. λ eli may be evaluated from bond length and bond angle variations, as is its homogeneous analog in Equation 1.67. For the evaluation of λ elo , the same model as in the homogeneous case is used, yet the electrode is replaced by the electrical image of reactant for the initial system (or that of the product for the final system) through the electrode, which acts as an electrostatic mirror. However, since these images are imaginary, only the real contributions are considered for the final evaluation of λ elo , given as follows: λ elo = N A

e2  1 1  1 1 2 −  − n  8π  εop εs   aR rΦ 

(1.106)

where aR is the equivalent radius of R (or P) rΦ is the distance between the centers of the R (or P) molecule and its electrical image when R (or P) is at the electron transfer site In practice, Marcus approximates rΦ by 2aR, assuming that the molecule is in contact with the electrode surface when at the electron transfer site. Thus, it follows that λ elo is given in Equation 1.107 [43a,b]:

(λ ) el o

Marcus

= NA

e2  1 1 1 2 −  n  8π  εop εs  2aR

(1.107)

Hush [78c] disagrees with the formulation in Equation 1.106, considering that the solvent polarization does not affect the potential in the double layer, the compact layer acting as an electrostatic Faraday cage. Thus, the incidental solvent fluctuations energetically affect only R or P, without consideration of their images (note that this amounts to considering rΦ, as infinite in Equation 1.106). Thus, Hush proposed the expression in Equation 1.108 for the outer shell reorganization energy:

(λ ) el o

Hush

= NA

e2  1 1 1 −  n2  8π  εop εs  aR

(1.108)

Interestingly, the Marcus heterogeneous reorganization energy (λ elo )Marcus in Equation 1.107 is exactly half that, λ oRR, of the homogeneous isotopic electron transfer: R* + P ⇌ P* + R whereas that proposed by Hush is identical to λ oRR. Based on this analogy, an elegant series of experiments by Kojima and Bard [80] led to the conclusion that Hush’s reorganization energy is more adequate [81].

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Introduction of λ el = λ elo + λ eli in Equation 1.105, when taking into account Equations 1.98, 1.101, 1.102, and 1.104, allows the expression of the current to be predicted as in Equation 1.109, where k(E) is the apparent rate constant for an electrode potential E. k(E) is given in Equation 1.110 under a formulation modeled on that of the empirical Volmer and Butler law [82], where the transfer coefficient α is given in Equation 1.111. i = nFAk ( E )[(R )0 − (P )0 e nF ( E − E

0

k ( E ) = kt0e( αn− Z R ) FΦs /RT e − αnF ( E − E

α=

) /RT

0

]

(1.109) (1.110)

) /RT

1 F (E − E 0 − Φs ) + 2 4λ el

(1.111)

The true standard heterogeneous rate constant kt0 is expressed in Equation 1.112 as a function of λel. Yet note that usually an intrinsic standard heterogeneous rate constant k0 in Equation 1.113 is defined and used rather than kt0. kt0 = Z ele −λ

el

(1.112)

/ 4 RT

(1.113)

k 0 = kt0e( αn− Z R ) FΦs /RT

Provided the exponential term in Equation 1.113, called the Frumkin correction, is invariant with E, that is, when Φs does not depend on the electrode potential, both are constant and characterize the R/P electron transfer independently of the electrode potential. Otherwise, kt0 is preferred. In the following, we nevertheless use k0 for simplicity in the formulations. 2. Transfer Coefficient α The formulation in Equation 1.110 shows that the apparent heterogeneous rate constant k(E) depends on the potential. This dependence has been observed experimentally since the earliest times of electrode kinetics and is at the origin of the so-called Tafel plots [83]. Indeed, when the potential E is sufficiently different from E0, for example, nF(E − E0) ≪ 0, the term relative to (P)0 in Equation 1.109 vanishes, and the current is given by the following: i = nFAk ( E )(R )0 = nFAk 0e − αnF ( E − E

0

) /RT

( R )0

(1.114)

When (R)0 is maintained constant, then a plot of the logarithm of the current versus the potential yields a straight line of slope αnF/RT, provided that α is constant. Thus, such observations led Butler and Volmer, two pioneers in this area, to propose an empirical law [82], named after them, in the form of Equation 1.109, with α being in their case an adjustable constant parameter. α is called the transfer coefficient after the following observation: when the driving force, that is, −nF(E − E 0) in Equation 1.114, is increased, only a fraction α of the additional energy is used to increase the apparent rate constant. Thus, α is the fraction of the driving force transferred for kinetic purposes and plays a role identical to that of the Brönsted coefficient in organic chemistry (see Section II.C.6). From Equation 1.111, it is seen that, provided that λ el term is large enough, α remains constant with the potential and is close to 0.5. This is generally in agreement with observations dealing with reactions with low values of kt0 (compare the λ el term in Equation 1.112). However, for organic molecules with small inner shell contributions and rather large radii, λ el (≈ λ elo ) is not expected to be

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Basic Concepts

log k(E)

0

–2

α

–4

–6

0.5

0.2 0.25

(a)

(E – E0), V

0.25

0.5

–0.25

–0.25

–0.5

(E – E0), V

(b)

FIgURE 1.16 (a) Variations of the apparent heterogeneous rate constant k(E) in Equation 1.110 with the electrode potential for the reduction of (CH 3 )3 CNO2 in DMF at the hanging mercury drop. A dashed line with slope 0.5 is positioned at E = E 0 to emphasize the curvature. (b) Resulting variations in the transfer coefficient α in Equation 1.111. (Experimental data from Savéant, J.-M. and Tessier, D., J. Phys. Chem., 81, 2192, 1977.)

extremely large, and therefore variations of α with the potential should be observed. This is indeed the case in practice [84], as illustrated by the curvature of the logarithmic plot of k(E) with the potential in Figure 1.16 for the reduction of 2-methyl-2-nitropropane, in DMF at the HMDE [84]. −

(CH 3 )3 CNO2 + e ⇌ (CH 3 )3 CNO•2

(1.115)

Moreover, the plot of the corresponding α(E) variations in Figure 1.16b shows that α depends linearly on the potential with a slope (0.249 V–1), in close agreement with that predicted (0.210 V–1) from the Marcus reorganization energy (29 kcal/mol) determined from Equation 1.107. 3.

Reversible and Irreversible Electron Transfers: Their Role in the Meaning of Oxidation or Reduction Potentials From the current expression in Equation 1.109, it is seen that when k0 is extremely large, the bracketed term on the far right must tend to zero in order that the current remains finite. Thus, in these conditions,* (R ) x =0 ≈ (P ) x =0 e nF ( E − E

0

)/ RT

which may be written under the form of the classic Nernst law in Equation 1.116. Thus, provided the heterogeneous rate constant is large (>0.1 cm/s), the electron transfer is controlled entirely by thermodynamics, and the relative activities of the oxidant and the reductant at the electrode surface obey the Nernst law in Equation 1.116. The electron transfer is thus termed reversible or nernstian. E = E0 +

RT (R ) x =0 ln nF (P ) x =0

(1.116)

In other words, when (J)x = 0 = (J)bulk, with J = R and P, because of a large excess of material at the electrode surface, for example, the electrode potential is given in Equation 1.116 independently of the current flow. Such an electrode is then suitable for use as a reference electrode (large k0, large surface area A, and excess of material). * Hereafter, (R)0 and (P)0 are considered to be the activities at the electrode surface and denoted (R)x = 0 and (P)x = 0 to recall this definition, x being the distance from the electrode. Yet, this usual notation is somewhat abusive since x = 0 does not imply a close contact but rather that R and P are located at the end of the diffuse layer (compare Scheme 1.5 and the associated discussion).

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Conversely, when k0 is extremely small, the current flow is negligible in the potential range around E0. Thus, in order that a measurable current may flow, the value of E − E0 must be significantly increased (compare k(E) in Equation 1.110). For example, in order that a cathodic current (n > 0) be observed, E must be considerably more negative than E 0. Thus, in Equation 1.109, 0 e nF ( E − E ) /RT → 0 in the potential range where the wave is observed, and Equation 1.109 can be rewritten as in Equation 1.117: i = nFAk 0e − αnF ( E − E

0

) /RT

( R ) x =0

(1.117)

Similarly, the observation of an anodic current requires that E is considerably positive to E 0. Then, in 0 the potential range where the reverse anodic current is observed (n > 0 in this case), e nF ( E − E ) /RT → ∞, and Equation 1.109 becomes i = −nFAk 0e(1−α ) nF ( E − E

0

) /RT

( P ) x =0

(1.118)

In these situations, the current flow is entirely controlled by the kinetics of the heterogeneous electron transfer, which is then said to be slow or irreversible. The current–potential characteristic of the redox couple under study then reflects thermodynamic (E 0) and kinetic (k 0) properties. Indeed, the driving force E − E 0 is used to overcome not only thermodynamic but also kinetic limitations. In this presentation, we discussed the reversible or irreversible nature of the electron transfer on an absolute basis. Yet, in practice, the notion of measurable current flow is obviously related to a given scale. In electrochemistry, the usual scale referred to is based on the mass transfer rate. Thus, for a given mass transfer rate, a given current density is expected to flow through the electrode. This limit is defined by the current density that would be observed if the electron transfer under investigation is nernstian. Within this context, it is easily understood that any electron transfer is nernstian at sufficiently small mass transfer rates. Similarly, it should become slow or irreversible for large mass transfer rates, because then, the current scale is made very high. In practice, the largest mass transfer rates routinely achievable with standard equipment are such that if k0 > 0.1 cm/s, the electron transfer may be considered intrinsically nernstian. A second point not considered in the discussion is related to the possible involvement of chemical reactions. For the sake of simplicity, let us restrict the presentation to the most common situation in which the product P formed upon electron transfer is consumed by a fast chemical step, as outlined in Scheme 1.6. It is easily seen from this simple representation that as soon as the rate of the chemical removal of P is larger than that of the back electron transfer, the RDS of the overall process is the forward electron transfer. Thus, independent of the intrinsic value of k0, the Butler–Volmer law in Equation 1.109 simplifies to that in Equation 1.117, because (P)x = 0 is made negligible. As a result, the R/P electron transfer presents all the kinetic characteristics of a slow electron transfer [85]. However, since this behavior is not related to the intrinsic value of k0, the current–potential curve may be observed in the close vicinity of E0 or even positive to E0 for a reduction (negative to E 0 for an oxidation) [81]. This is an important notion to recall when trying to relate this position of an irreversible wave to thermodynamics. In absence of additional knowledge E 0 may be located on either side of the observed wave. kf P

R + ne kb

SCHEME 1.6 Competition between forward and backward electron transfer reactions and follow-up reaction at the level of the electron transfer intermediate product P.

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55

D. ADSORPTION PHENOMENA Many solutes or follow-up products of electrochemical reactions may have a large tendency to adsorb on electrode surfaces [86]. This may be related to long-range electrostatic interactions, such as those discussed in the presentation of the double layer; there is then no specificity of the ions attached, particularly as concerns their chemical properties. In such cases, the phenomenon is said to be nonspecific and is akin to nonspecific salt effects or ion pairing in organic chemistry. On the other hand, because of particular chemical affinities between the electrode surface and a substrate, specific adsorption occurs, which involves generally strong interactions and may result in the creation of organic layers at the electrode surface. Analogies of the latter are given by complexation reactions or specific salt effects (e.g., Li + with an enolate ion or Bu− Li + solutions) in organic chemistry or chemisorption and physisorption in heterogeneous catalysis. Both classes are distinct because of the nature of the electrode–adsorbed species interactions. Yet from a phenomenological point of view, a better classification seems to be that based on the possible electroactivity or conductive properties of the adsorbed material. When adsorbed, electroinactive species may give rise to partial or total blocking of the electrode surface, which may then not be able to transfer electrons to or from other electroactive molecules. Thus, depending on the nature of the electrode coverage, various phenomena may be observed, ranging from an apparent decrease in the electrode surface (large and separated islands of adsorbed species) to an apparent decrease in the electron transfer rate but with the appearance of unchanged electrode surface area (micrometric islands at micrometric distances from each other) [87]. Related also to this class is the natural blocking of electrodes by oxide layers or functional groups at carbon electrodes [88]. Conducting (via electrons or via permeation of electroactive species) adsorbed species usually introduce small disturbances in electrochemical behavior compared with the former class. Yet obviously, the diffuse layer is extremely affected vis-à-vis the bare electrode. Thus, the rate constants k0 may be modified to a high degree because of large changes in the Frumkin correction.* Such effects may easily explain, at least on a qualitative basis, the well-known dependence of k0 on the size of the supporting electrolyte cation (reductions) or anion (oxidations), as well as on ionic additives [89]. The third class involves electroactive adsorbed species, yet it also encompasses the case in which, although originally inactive, the species is activated or modified in the adsorption process to give rise to the formation of electroactive adsorbed molecules. Obviously, all these effects discussed for the two previous classes are observed for this class, but additional effects are observed for the adsorbed species itself. Indeed, adsorption modifies (1) the concentration of the species at the electrode surface vis-à-vis that predicted from its bulk concentration, and (2) the energetics of the electron transfer reaction. The latter point arises because of additional thermodynamic contributions to be considered in the R0/RΦ or P0/PΦ equilibria in Scheme 1.5 and also from changes in the reorganization energies. The former point leads to a variety of behaviors, since it is obviously a function of the kinetics of adsorption–desorption phenomena [90]. A third additional effect of electroactive adsorbed species arises from their possible ability to act as electron transfer mediators for other species that are reduced or oxidized with more difficulty. This special class (the term electrode modification is then more adequate) is at present a very active area of research.

E.

COUPLED CHEMICAL REACTIONS

1. Electrochemical versus Homogeneous Chemical Reactivities In the preceding sections of this chapter, much attention was devoted to the essential act of electron transfer. Yet although being, for fundamental or analytic purposes, the subject of a large body of the electrochemical literature, pure electron transfer reactions generally are of less interest in terms of * Note that for n = 1, an increase by 0.2 V in ΦS results in the multiplication of k0 by approximately 50 when α = 0.5, and by approximately 350 when α = 0.75, as determined from Equation 1.113.

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preparative electrochemistry. In the opinion of this author, electron transfer at an electrode must be considered as a particular class of activation of molecules to enhance their chemical reactivity. As such, electrode kinetics must be understood (as for other methods of chemical activation) in order to control and eventually direct the overall process to the selected target (compare Chapter 10). From the preceding discussion, it should appear that electrochemistry affords a very facile and precise way to generate highly energetic intermediates via control of the electrode potential. Note in this respect that one electron exchanged over a potential difference of 1 V amounts to injecting 1 eV in a molecule, that is, approximately 23 kcal/mol. Owing to the possibility of performing electrochemistry over potential differences of several volts, it is easily seen that the method may involve energies comparable to those of most chemical bonds and largely beyond most activation energies. Thus, highly energetic intermediates may be generated under mild and precisely controlled conditions. Another useful aspect of electrochemical generation of chemical intermediates is related to the current flow through the cell. Indeed, the current is a direct measure of the production rate of the intermediate. It ensures that the production rate, an important parameter in product selectivity, is easily controlled and adjusted. At this point, two questions must be addressed: Are these intermediates identical to those obtained under usual chemical conditions? Do they react under conditions matching those of homogeneous experiments? The answer to the first question is obviously positive. Indeed, since the substrates are generally neutral, the primary electron transfer affords paramagnetic intermediates (anion or cation radicals) that may undergo selected cascades of reactions to afford most of the well-identified intermediates of organic chemistry: radicals, carbocations, carbanions, carbenes, and so on. Neutral radicals may also be generated by reduction or oxidation of stable cations or anions, respectively. It is thus seen that all the basic intermediates of organic chemistry may be electrochemically generated, provided an electroactive precursor exists. Yet this is not always a serious limitation, since the sought intermediates may be eventually formed indirectly, that is, through an electroactive proreagent, as discussed elsewhere in this book (Chapters 15, 43, and 44). The second question is more difficult to answer. Indeed, since the electrical perturbation extends only over a few angstroms from the electrode surface, the medium in which the electrogenerated intermediates react is identical to the bulk of the solution, which favors a positive answer. On the other hand, they react under essentially nonisotropic conditions because of the existence of concentration profiles, that is, of concentration changes with distance from the electrode (see Figure 1.15). Yet this is a situation identical to that encountered when heterogeneous reagents or polyphasic conditions are used in usual homogeneous chemistry. A more serious difference is related to the nature of the media often used under electrochemical conditions. As explained in Section III.A.3, electrochemical solutions must generally be rather conducting, a requirement met by using solvents with rather high dielectric constants and an excess of inert electrolyte (see Chapter 7). Thus, although solvents, such as liquid ammonia, or media with smaller dielectric constants, such as methylene chloride, THF, DME, or even arenes and alkanes, can be used [91,92], the preference of most organic electrochemists is for convenience or by tradition, to such solvents as acetonitrile and dimethylformamide. The latter are prone to intervene in the reaction paths of the electrogenerated intermediates via their nucleophilic or acidobasic properties, for example, as well as via hydrogen atom transfers to radicals [93]. The same is true for the so-called inert electrolytes, frequently quaternary ammonium salts, known for their acidic (Hofmann reaction) abilities, as well as for being relatively good hydrogen atom donors. Thus, the usual electrochemical media may induce reaction paths that are quite different from those observed with identical intermediates under conventional organic conditions. We want to emphasize, however, that this is not to be viewed as a limitation of the method. Indeed, these induced reaction paths may be used with great profit, as in organic chemistry when methanol is used as the solvent to incorporate methoxy groups. Several examples of this strategy are given in various places in this book. On the other hand, as mentioned earlier, a large variety of solvent–electrolyte systems may be used in electrochemistry [71,91], although the general body of reported data pertains to a limited class of media.

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2. Chemical Reversibility and Irreversibility: Their Role in the Meaning of Oxidation or Reduction Potentials When considering a simple electron transfer reaction at an electrode, as in Equation 1.119, one can conceive easily that there are, a priori, four limiting situations in which chemical reactivity may affect the electrochemical measurements: R may be consumed or produced, and P may be consumed or produced. R + ne ⇌ P (E 0 )

(1.119)

In practice, all four possibilities may be combined, which leads to a large variety of electrochemical mechanisms. Thus, our purpose in this section cannot be exhaustive if we wish to keep some generality. We want, however, to describe qualitatively the role and the effect of chemical reactions on the overall electrode kinetics. Thus, for the sake of simplicity, we discuss the case of simple EC (electrochemical–chemical steps; see Section III.E.3) sequence, in Equations 1.120 and 1.121 as a typical example: R + ne ⇌ P( E 0 , k 0 , α)

(1.120)

P → ⋯(k )

(1.121)

The chemical step in Equation 1.121 is homogeneous and thus intervenes while P, which is formed at the electrode, is transported to the bulk of the solution via the mass transfer mechanisms discussed earlier. From this simple consideration, it is easily understood that an important factor is the relative magnitude of the rate constant of the chemical reaction vis-à-vis that of mass transfer, as outlined in Scheme 1.7. Obviously, when the mass transfer rate is much larger than that of the chemical steps, the transfer of P from the electrode surface to the bulk solution is not affected by the presence of kinetics. Thus, the electrochemical behavior of the EC sequence in Equations 1.120 and 1.121 is identical to that of a simple electron transfer mechanism. The R/P electron transfer is then said to be chemically reversible but may be nernstian or slow, according to the magnitude of k0 in Equation 1.120.* In the converse situation, three limiting cases may be observed. In the first case, the electron transfer is intrinsically slow. The RDS of the overall process is then the forward electron transfer, and the redox system is said to be slow and chemically irreversible. The electrochemical wave is then observed at potentials sufficiently different from E 0 for n(E − E 0) ≪ 0 (see Section III.C.3). The two other situations are encountered when k0 is large enough for the electron transfer to be nernstian in the absence of the follow-up chemical step. Mass transfer Rx = 0 + ne

kf kb

Px = 0 k

Chemical sequence

SCHEME 1.7 Competition between forward and backward electron transfer reactions, diffusion, and followup reaction at the level of the electron transfer intermediate product P. * To avoid possible confusion with the reversible or irreversible nature of the electron transfer step per se, we suggest that reversibility or irreversibility be reserved to the chemical part and the electron transfer be designated by nernstian or slow, according to its rate versus the mass transfer one. In the following, we use this terminology.

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When the chemical reaction is slow compared with the backward electron transfer, the reduction or oxidation in Equation 1.120 remains at equilibrium. The electrode potential is then given by the Nernst law: E = E0 +

RT (R ) x =0 ln nF (P ) x =0

(1.122)

Yet because of the chemical reaction, the activity of P at the electrode surface is considerably decreased when the chemical rate exceeds that of the mass transfer. From Equation 1.122, it is seen that the electrode potential is then positive (for a reduction; negative for an oxidation) to that observed for the same current density, but in the absence of the follow-up reaction. As a result, the current–potential characteristic is observed in a potential range positive to E 0 for a reduction and negative to E 0 for an oxidation. The system is then said to be nernstian and chemically irreversible. In the last situation, the rate of back electron transfer, although intrinsically fast, is slower than that of the chemical step. Thus, as in the first situation, the RDS is the forward electron transfer, and the redox system is said to be slow and chemically irreversible. Yet since k0 is large, the electrochemical wave is observed before E 0 [81,85]. From this description, it is seen that, depending on the exact degree of competition between the four rates in Scheme 1.7, the electrochemical wave is observed in potential ranges that may be positive or negative to E0, as summarized in Table 1.4. Without a quantitative treatment of the pertinent electrochemical data, it is almost impossible to decide the position of E0 vis-à-vis the potential location EW, where the wave is observed. This is an important caveat to remember when using published potential values, such as peak potentials or half-wave potentials, instead of E0 values. The experimental difficulty is even more severe in practice. Indeed, when one considers a series of related chemicals, there are great chances that E 0, k0, and k vary uniformly with respect to each other because they relate intimately to the orbital energy and characteristics of the acquired (LUMO) or lost (HOMO) electron [16,29]. Thus, the potential EW at which the electrochemical wave is observed may correlate with E0, since it depends on the three figures. Such a case is presented in Figure 1.17 for the series of alkylbenzenes [29] already mentioned in this chapter, in which the peak potentials of the chemically irreversible voltammograms (at 0.1 V/s) correlate with E0 with a slope close to unity. To conclude this section, we discuss the special case of follow-up equilibria, that is, when the chemical step in Equation 1.121 is reversible [93]. R + ne ⇌ P ⇌ Z ( E 0 , k ( E ); kf , kb )

TAbLE 1.4 Nature and Locationa of the Electrochemical Wave Observed for an EC Sequenceb as a Function of the Mass Transfer Rate k0 versus Mass Transfer Rate k versus Mass Transfer Rate Small Large

a

b c

Small

Large

Slow reversible,c n(EW − E0) < 0 Slow irreversible,c n(EW − E0) < 0

Nernstian, reversible,c EW ≈ E0 Nernstian or slow irreversible,c n(EW − E0) > 0

EW, any characteristic potential of the wave: E1/2, Ep, Ep/2,…, E0, standard reduction potential of the electron transfer step, involving an exchange of n electrons (n > 0 for a reduction; n < 0 for an oxidation). k0, intrinsic standard heterogeneous rate constant; k, rate constant of the homogeneous follow-up chemical reaction. Chemically reversible or irreversible.

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E 0 , V vs NHE

2.5

2.0

2.0

2.5

Ep , V vs NHE

FIgURE 1.17 Correlation between E 0 and anodic peak potentials Ep of the chemically irreversible cyclic voltammograms at v = 0.1 V/s determined for an extended series of alkylbenzenes. (Data from Howell, J.O. et al., J. Am. Chem. Soc., 106, 3968, 1984.)

Obviously, when K = kf/k b ≪ 1, the equilibrium lies toward P and does not affect the R/P electron transfer. In the converse situation, and when kf and k b are larger than the mass transfer rate, a rapid equilibrium displaced toward Z establishes, and (P)x = 0 ≈ (Z)x = 0/K. Introduction of this relation in the current rate law Equation 1.109 then yields Equation 1.123, which is identical in its formulation to a Butler–Volmer law but for a hypothetical electron transfer R + ne ⇌ Z whose formal potential is E 0′ = E 0 + ( RT /nF ) ln K , as evidenced by the formulation in Equation 1.124. 0  e nF ( E − E ) /RT  i = nFAk ( E ) (R ) x =0 − ( Z) x =0  K  

i = nFAk ( E )[(R ) x =0 − ( Z) x =0 e nF ( E − E

0′

)/ RT

]

(1.123)

(1.124)

This constitutes the basis of the operational significance of formal potentials introduced in Section II.B.5. Yet the formal analogy between Equations 1.109 and 1.124 should not mask the fact that k(E), that is, α, kt0, or k0 in Equation 1.124, remains related to the real electron transfer R + ne ⇌ P, that is, to E 0, not to E 0′ (compare Equations 1.110 through 1.113). 3. Classification of Coupled Chemical Reactions From this discussion, it is seen that the electrochemical behavior of a given electron transfer reaction may be extremely perturbed by the intervention of coupled chemical steps. The resulting modifications obviously depend on the exact nature of the chemical sequence associated with the electron transfer step. It is thus convenient to classify the different possible reaction schemes that may arise. Yet as in organic chemistry, for the basic mechanisms (e.g., SN1, SN2, E1, and E2), such a classification needs to be viewed only as a general frame to which real mechanisms are referred. Indeed, the latter may often be more subtle (compare, e.g., SNi or E1cb vis-à-vis the earlier mechanisms) or involve cooperative or competitive combinations of the elementary schemes presented later. A  second

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similarity with the organic basic schemes is that there is no unity in the denominations of the various sequences of electron transfers and chemical steps. Indeed, each school or laboratory has proposed its own nomenclature, which has received a different audience and may have evolved with time. In the following, we use as much as possible the most frequent designations, and when not possible, we refer to the nomenclature used in Savéant’s group, owing to its large contributions to this area [94]. The following classification is based upon the fact that electrochemical behavior depends on the location of the electron transfer step in the overall sequence. Also, electron transfer steps are considered to involve the exchange of a single electron, that is, n = ±1 (n = 1 for a reduction and n = −1 for an oxidation), a normal situation in organic electrochemistry. Indeed, although overall stoichiometrics involving the net exchange of many electrons are frequently observed, the energetic and activation requirements for a direct polyelectron transfer are too high for any to exist concurrently with a rapid succession of single-electron transfer and chemical steps. Within this frame, the most frequently encountered situations are the following.* a.

Preceding Chemical Reaction: CE Mechanism  k  Z ⇌ A  kf , kb , K = f  k b  

(C)

A + ne ⇌ B ( E 0 )

(E)

The electroactive species A is generated by a reaction, generally an equilibrium displaced toward Z, that precedes the electron transfer step. Z is the reactant introduced in the cell or the predominant form of the reactant in the reaction medium. These reaction schemes were introduced to rationalize the various electrochemical phenomena observed during the reduction of certain aldehydes in aqueous solutions. Indeed, in water, formaldehyde, π-electron-poor heterocyclic aldehydes, and a few aldehydes with strongly electron-attracting groups exist as their nonreducible hydrated form in rapid equilibrium with the reducible carbonyl (Scheme 1.8). When the rate of the Z ⇌ A interconversion is slow vis-à-vis the rate of mass transfer,† it acts as the RDS for the overall sequence. The currents observed are then abnormally low for the mass transfer rate and reactant concentration considered and are said to be kinetic currents. Conversely, when both rates are fast, the preceding chemical reaction acts as a rapid preequilibrium, and (A) = K(Z). The electrochemical data are then formally identical to that of the virtual Z + ne ⇌ B electron transfer, with a formal potential E 0′ = E 0 −

RT K ln nF K + 1

OH RCHO + H2O

RCH OH

R C

RCHO + e

O–

H

SCHEME 1.8 Reduction mechanism of a hydrated carbonyl species (CE mechanism). * Hereafter, we use the following notations: E stands for heterogeneous electron transfer, C for a chemical step. In chemical reactions, A, B, C, and so on designate the electroactive molecules or their reaction product(s), whereas Y, Z, and so on indicate coreagents or nonelectroactive species. † In the following, rapid, fast, and slow always refer to a comparison to the rate of mass transfer.

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Basic Concepts B

A + ne

B+Z

SCHEME 1.9

(E)

( n = ±1)

(C')

etc

A+Y

Schematic representation of redox catalysis (ECʹ mechanism).

b. Redox Catalysis or Mediated Electron Transfer: EC′ Mechanism In this sequence, the A/B couple acts as an electron transfer mediator, which allows the reduction or the oxidation of the electroinactive (at the A/B wave potential) Z species (Scheme 1.9). The electroactive species A is regenerated via the homogeneous electron transfer step C′, which results in a considerable enhancement of the current observed for the A/B wave as soon as the C′ step is rapid. Various subclassifications exist according to the exact nature of the chemical step, which may eventually be a succession of elementary steps with formation of intermediate products. As explained earlier for the displacement of endergonic electron transfer steps, the C′ step occurs because it is continuously pulled to the right by the further chemical reaction of the Y species. Note that most of this book is devoted to this class of mechanisms. c.

Consecutive Reaction: EC Mechanism (E)

A + ne ⇌ B (n = ±1) B→C

(C)

This sequence is one of the most ubiquitous in organic or organometallic chemistry and has already been discussed in preceding parts of this section. The chemical step may involve reactions (or a succession of chemical reactions) of high-order modularities, not only first-order reactions as suggested by the notation. Classic examples are given by activated olefin electrohydrodimerization mechanisms (Scheme 1.10)*. d. Multielectron Reactions: ECE and Related Mechanisms As explained earlier, a direct exchange of many electrons at the same redox center appears to be extremely difficult because of energetic and activation requirements. From the formulation in Equation 1.106, it is indeed noted that the exchange of 2e involves a reorganization energy λ elo four times larger than that for a single-electron exchange at the same center. For this reason and because of the free energy relationship in Equation 1.105, a two-electron step without formation of an intermediate should correspond to a slower process than the corresponding EE sequence as follows: −

(1.125)

A + 2e ⇌ A • + e ⇌ A 2 − H

H CH2

+e

C

CH2

– C Z

Z

H 2 CH2

H

H – C

– C Z

Z

(CH2)2

– C

2H+

Z

(CH2)4

Z

Z

SCHEME 1.10 Radical anion–radical anion coupling sequence during the electrodimerization of activated olefins. * We use here only one example called radical–radical coupling though many other coupling steps may occur [94].

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On the other hand, in the second step of this sequence, the electron is transferred to a molecule with an additional negative charge compared with A. Thus, at least a coulombic repulsion term must be overcome, which implies that the second electron transfer is necessarily performed at a more negative potential than the first (by at least 0.250 V for an equivalent radius of 2 Å for the redox centers and, e.g., ε ≈ 37.5 for acetonitrile as the solvent). As a result, an overall EE sequence should not be observed at the potential of the A reduction but should occur in two separate waves (see Chapter 11). However, there are two ways in which this conclusion may be invalid, and both need the relay of a chemical reaction. The first possibility is that the reorganization energy associated with the first electron transfer is considerably larger than that of the second. Such a difference is obviously not related to the outer shell components, which are necessarily close, as evidenced by Equation 1.106. Thus, the origin for the differences must be sought in the inner shell reorganization energy, which implies that a profound molecular reorganization must be involved during the first electron transfer. Examples of this situation are given by the numerous cases in which two redox centers are electronically coupled in the same molecule, as in O2 N −(CH 2 )n − NO2. Similarly, for inner sphere electron transfers, the molecular moiety bearing the negative charge may be lost, leaving a neutral and easily reducible species. This is so in the non-concerted reduction of alkyl halides to alkanes: RX + e → R • + X

(E1)

R• + e → R −

(E 2)

R − + H + → RH

(C)

in which the first electron transfer leads to the carbon–halogen bond cleavage and affords an easily reducible neutral radical [95]. However, for most organic electroactive molecules, such large reorganization energies cannot be invoked, owing to the extended delocalization of the orbitals participating in electron transfers. Yet the interposition of a chemical step (possibly a simple conformational change) between the two electron transfers results in the same effect, as soon as the chemical reaction gives rise to a species more easily reducible (or oxidizable, in oxidations) than the starting reactant. A + ne ⇌ B E10

(E)

B→C

(C)

C + ne ⇌ D

(E ;n(E 0 2

0 2

) )

− E10 > 0

(E)

The intermediate formed is then reduced (or oxidized) at the electrode or in the solution owing to the large exergonicity of the resulting reaction: B+C → A + D

( ∆G

0

(

)

= −nF E20 − E10 ≪ 0

)

(DISP)

In the former case (second electron transfer at the electrode), the overall sequence is termed ECE; it is designated a DISP (rather than an ECDisp) mechanism in the second case because the homogeneous electron transfer step is akin a disproportionation reaction. Indeed, the latter involves an electron transfer between two chemically related molecules with the same oxidation numbers. In practice, the distinction between the two mechanisms is nearly impossible on experimental grounds [96] because their RDS, the B → C reaction, is identical.

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Basic Concepts

A typical example of this sequence is given by the non-concerted reduction of aromatic halides: −

ArX + e ⇌ ArX • −

ArX • → Ar • + X − Ar • + e → Ar −

−

and/or Ar • + ArX • → Ar − + ArX Ar − + H + → ArH

As in the EC electrochemical sequence, a large variety of mechanisms belong to the general ECE– DISP frame and differ according to the exact molecularity of the interposed chemical step, which may also be a succession of elementary steps. Similarly, cascades of ECECE… sequences are possible, as in the classic reduction of nitro derivatives to hydroxylamines, which involves four electrons and four protons [97]: ArNO2 + 4e + 4ROH → ArNHOH + 4RO − + H 2O

(1.126)

e. Zero-Electron Reactions: Electron Transfer Catalysis of Chemical Reactions In the preceding ECE sequence, both electron transfers are of the same kind, that is, both reductions or oxidations, and thus their combination gives an increase in the absolute value of the number of electrons exchanged. Yet when the second electron transfer is of a nature opposite to the first, the overall number of electrons consumed is zero. Thus, the sequence amounts to an electron transfer catalysis of a nonredox chemical reaction [98], as outlined in the following sequence: A + ne ⇌ B

(E)

B→C

(C)

C − ne → D and/or C + A → D + B

(E)

Naturally, such mechanisms are observed only when the overall A→D reaction is exergonic but is considerably slower than that occurring at the level of the reduced or oxidized states B and C. These processes have been shown to be general and to encompass a large variety of chemical reactions. Among the most documented is the SRN1 reaction, which amounts to an aromatic nucleophilic substitution. Although initially postulated by Bunnett [99a], the confirmation and the subtleties of its mechanism have been established on electrochemical grounds [100]: −

ArX + e → ArX • −

ArX • → Ar • + X − −

Ar • + Nu − → ArNu• −

−

−

ArNu• − e → ArNu and/or ArNu• + ArX → ArNu + ArX • Overall ArX + Nu − → ArNu + X −

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4. Electrochemical Reaction Mechanisms and the Principle of Microscopic Reversibility As understood from the large variety of possible reaction sequences outlined in Section III.E.3, a large number of reaction paths may be invoked a priori to explain the same overall electrochemical transformation. Yet owing to the great capabilities of electrochemical techniques for the identification of reaction mechanisms, the problem of their distinction frequently arises. For example, let us consider the very common two-electron two-proton electrochemical sequence as follows: A + 2e + 2H + ⇌ AH 2

(1.127)

The latter is involved, for example, in the reduction of aromatics to their dihydro derivatives, of quinones to hydroquinones, carbonyls to alcohols. A priori, there are six plausible different possibilities of transferring two electrons and two protons, depending on the exact order of the steps. These six paths + − − • involve the participation of seven different intermediates: A • , A2−, AH • , AH • , AH + , AH 2+ 2 , and AH 2 . It is thus seen that the discussion of the possible mechanism(s) may rapidly become an overwhelming task, particularly when more steps are involved, as in the 4e + 4H + sequence in Equation 1.126. Thus, in order to discuss the exact route followed, a convenient representation of the different possible schemes is highly desirable. Such a representation of all possible paths is given by square-schemes diagrams [101], such as that shown in Scheme 1.11 for the A/AH2 reduction. The horizontal transformations correspond to electron exchanges, whereas the vertical transformations involve a proton transfer. Based on chemical grounds, several routes may be ruled out in Scheme 1.11. For example, starting from an aromatic hydrocarbon A, routes δ, ε, and φ may be excluded, except in a superacidic medium. + This is also certainly true for the γ route, which involves a difficult protonation of AH • to AH 2• . Thus, only two routes, α and β, appear chemically reasonable under usual electrochemical conditions. In practice, both may be followed, path α prevailing under weakly acidic media and low cathodic potentials, and path β predominating in less acidic conditions and at more negative values of the potential. Note, however, that additional complications, not considered in Scheme 1.11, involve the possible role of homoge− neous electron exchanges. Thus, AH • may be reduced at the electrode but also by the strong reductant A • . The reverse transformation may also be considered, as in the oxidation of aromatic dihydro derivatives to their parent hydrocarbon. Then, one starts from AH2, and routes α, β, and δ must be eliminated, except in very basic media. This is also certainly true for route γ, since the anion radi− cals A • usually have a marked basic character. Thus, under usual electrochemical conditions, only “Square-scheme”:

Possible mechanisms

A

A–

– A2

AH+

AH

AH–

A

A

(α)

AH2

A AH2+ 2

+

AH2

(β)

A

AH2

(γ)

AH2

A

A

AH2

(δ)

AH2

AH2 (ε)

AH2 (φ)

SCHEME 1.11 Square-scheme mechanism detailing all possible sequential electron transfer and chemical steps that may be involved in the two-electron reduction of species such as quinones. Left: global squarescheme; Right: representation of all individual possible pathways that may occur depending on the exact experimental conditions.

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Basic Concepts

routes ε and φ must be retained. Path ε predominates at less positive potentials and weakly basic conditions, whereas path φ should be observed at more positive potentials. Consider now a given Ar ⇌ ArH2 reaction under rather neutral conditions. It is inferred from the preceding discussion that route α is followed for the reduction, whereas the oxidation goes through path ε. At first glance, such a result seems to have to be rejected since it apparently contradicts the principle of microreversibility, which claims that both reactions should follow the same mechanism. However, both transformations do not occur at the same electrode potentials: the reduction Ar → ArH2 needs rather negative values, whereas the oxidation ArH2 → Ar is performed at rather positive values. Thus, the two reactions occur under driving forces of opposite direction, with an energetic difference of approximately 100–130 kcal/mol. Thus, invoking the principle of microreversibility under such conditions is no more reasonable than invoking it when comparing the Birch reduction to the oxidation of aromatic dihydro derivatives by a strong oxidant.

IV. MASS TRANSFER IN ELECTROCHEMISTRy A. FUNDAMENTAL ASPECTS OF MASS TRANSFER PROCESSES As mentioned in the preceding sections, mass transfer plays a crucial role in electrochemistry. The ubiquity of this phenomenon evidently arises because the electrons are exchanged at 2D surface boundary, but the reactants and products are dispersed in a 3D solution. Thus, in the absence of mass transfer, only a small layer (of a few molecular radii thicknesses) would exchange electrons with each electrode, as is the case, for example, when the electroactive material is adsorbed at the electrode (see Section III.D). As such, mass transfer is essential in controlling the success and rate of most electrochemical reaction. Another important feature of mass transfer processes is related to the very physical nature of the phenomenon. As such, it is easily quantifiable and predictable. Thus, the rate of mass transfer to and from an electrode may be determined a priori for a given electrochemical system. As a result, this rate may be used as natural built-in clock by which the rate of other electrochemical processes may be measured. Such a property was apparent in our earlier discussions related to electrode kinetics (electron transfer and coupled chemical reactions). Basically, it proceeds from the same idea as that frequently used in organic chemistry for relative rate constant determinations, when opposing a chemical reaction of known (or taken as the reference in a series of experiments) rate constant against a chemical reaction whose rate constant (or relative rate constant) is to be determined. Many such examples exist in the organic literature, among which are the famous radical-clocks (Scheme 1.12). However, because of the intrinsic specificity of such chemical clocks, they cannot have a general use outside a given class of chemistry. Thus, a better analogy with the electrochemical approach involving mass transfer rates would be that in which a reaction rate constant is matched to the diffusion limit rate constant that represents the physical process of bringing two reactants to distances such that they can interact chemically (see Section II.C.1). Obviously, such methods k(

) Z

Z

I

(I)/[(II) (

)] = k/kclock

Z

kclock

II

SCHEME 1.12 Principle of using a product ratio I/II for evaluating the relative rates of two pathways involving a radical clock.

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Organic Electrochemistry

are not frequently used in homogeneous chemistry, principally because most chemical reactions, with maybe a few exceptions, among which are electron transfer or radical reactions, proceed with rates considerably slower than that of the diffusion limit. Another limitation to the development of such an approach is also related to the fact that diffusion limit is a rate of mass transfer (analogous to that involved in electrochemical experiments) but relates to mass transfer over atomic distances and is thus not prone to be adjusted at will. In this respect, electrochemical mass transfer is particularly convenient. As mentioned earlier, its physical nature and its necessary involvement in any electrochemical experiment confer on it one of the characteristics of the internal clock just described. On the other hand, that it is active in the vicinity of a macroscopic object, the electrode allows its precise control. In practice, mass transfer rates at electrodes are continuously adjustable from equivalent times of a few seconds to a few nanoseconds [102]. When this is coupled to the use of a diffusion limit in the nanosecond or less time scale, it explains why electrochemical transient techniques allow kinetic information to be obtained and rate constants determined for reactions whose lifetimes range from seconds to tens of picoseconds [103]. 1. Physical Processes of Mass Transfer: Fick’s First Law There are three physical processes involved in the transport of molecules in solutions: convection, migration, and diffusion. Convection may be the most intuitive and is present in all areas of chemistry; it corresponds to forced displacements of small volumic parts of the solution. The molecule is then carried within the fluid elements in which it is dissolved. The flux of particles Jconv through an elementary surface, that is, the number of moles of particles crossing a unit surface area per unit of time, is then simply expressed as J conv = Cv( x )

(1.128)

where v(x) is the velocity of the fluid normal to the surface C is the concentration of the particle in the fluid element It then follows that the mass transfer contribution is independent of the particle but is only a function of the local hydrodynamics (natural or thermal convection, forced convection, eddies, and so on). A large amount of work has already been devoted to quantifying the hydrodynamics of a solution [104]. Yet for our purposes here, it is sufficient to recall two conclusions that apply in most electrochemical situations. (1) Because of shocks and interpenetration of the fluid elements in motion, concentrations are nearly identical at any point of a solution far from any wall. (2) Near any wall or solid surface, solid–liquid frictions and bouncing phenomena result in the creation of an immobile solution layer in which no convective effects occur.* Obviously, the thickness δconv of this layer depends on the fluid velocity in the vicinity of the wall. Yet the resulting variation is small: in nonviscous solvents, thermal agitation and average vibrations result in δconv ≈ 100 μm, whereas magnetic stirring at maximum agitation results in δconv ≈ 10−20 μm. From an electrochemical point of view, it is easily inferred that the solution in a cell near an electrode is separable into two parts [105]: a stagnant layer adjacent to the electrode in which no convective motions occur and the remainder of the solution, which is homogeneous (bulk solution). Yet this is not a particularity of electrochemical methods since the same phenomena occur at any solid/ liquid interface, as when metal particles (e.g., reductions by Zn or Na) or any heterogeneous reagent is used in organic homogeneous chemistry, as well as in phase-transfer catalysis or related methods. * Note that the separation of the solution into two parts near a wall, that is, the immobile layer and the agitated homogeneous solution, is really simplistic since there is a continuous variation in the fluid velocity from the wall to the bulk [105]. Yet this dichotomous approximation is sufficient for most purposes and allows great simplifications in physicomathematical treatments of the transport problem.

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Basic Concepts

The other two contributions to mass transfer arise from forces exerted on the molecule considered. When the molecule or particle is charged and an external electrical field is applied, electrostatic forces are effective, and a positive molecule descends the electrical potential gradient, that is, goes in the direction of lower potentials, whereas a negative ion climbs the potential gradient. Because of frequent shocks and bouncing against solvent molecules, the molecule reaches a limiting velocity proportional to the potential gradient times its charge z, the proportionality coefficient γ depending on the molecules shape and size, as well as on the viscosity of the medium. This phenomenon, designated migration, then corresponds to a flux Jmigr given in the following: J migr = − γzF

∂Φ C ∂x

(1.129)

where x is the coordinate normal to the elementary surface crossed ∂Φ/∂x is the electrical potential gradient along the x axis C is the local molecule concentration Φ can have an external origin as in electrophoresis or can be the result of local changes in the distribution of other charges. In the presence of a sufficient excess of supporting electrolyte, the latter contribution is shown to be extremely small [106]. The third contribution is called diffusion, but a better name, although not used, would be “spreading.” Indeed, it corresponds to the natural tendency of nonrigid objects to spread so that the object’s constituents are leveled. This is well known in any field of chemistry and particularly disturbing in chromatography, since it is responsible for the enlarging of peaks with retention times. The physical origin of this phenomenon is easily understood when considering the chemical potential μ of a solute (for simplicity, concentrations and activities are assumed equal) given in Equation 1.130, where μ0 is the standard value. Thus, when C is not uniformly distributed along the x axis, one obtains the chemical potential gradient in Equation 1.131: µ = µ0 + RT ln C

(1.130)

∂µ ∂ ln C RT ∂C = RT = ∂x ∂x C ∂x

(1.131)

This potential gradient can be considered as a force that acts on the molecule, as occurs for a charged species submitted to an electrical potential gradient. Again this gives the molecule a limiting velocity oriented in the direction opposite to the gradient, that is, toward the lesser values of μ, and proportional to the force −∂μ/∂x. Because of the same collisions and bouncing effects as in the migration case, the proportionality coefficient is γ, identical to that in Equation 1.129. Thus, the diffusion flux Jdif is given as follows: J dif = − γ

∂µ ∂C C = − γRT ∂x ∂x

(1.132a)

In practice, one prefers to reformulate Equations 1.129 and 1.132a by introducing the diffusion coefficient D = γRT of the molecule in the medium considered, that is, J dif = − D which is commonly designated as Fick’s first law.

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∂C ∂x

(1.132b)

68

Organic Electrochemistry

In the most general situation, all three mass transfer modes cooperate algebraically; thus, the number of moles crossing a surface of unit area per unit of time is given as follows, which constitutes an extension of the Fick’s first law: J ( x ) = J conv + J migr + J dif = Cv( x ) − CD

zF ∂Φ ∂C −D RT ∂x ∂x

(1.133)

2. Relationships between the Electrochemical Current and Mass Transfers On the basis of the dichotomous representation of the solution near the electrode surface, Equation 1.133 simplifies in the two domains [105]. Within the stagnant layer and in the presence of an excess of supporting electrolyte, the two first terms are negligible, and one obtains Equation 1.134. Conversely, when x > δconv, the solution is macroscopically homogeneous; the diffusional contribution then vanishes, and the flux is given in Equation 1.135. 0 < x < δconv : J ( x ) = − D

∂C ∂x

(1.134)

zF ∂Φ   x > δconv : J ( x ) = C  v( x ) − D RT ∂x  

(1.135)

Equation 1.135 allows one to discuss more specifically the problem of ohmic drop. Indeed, the current density through any surface area A located at any point in the bulk solution is given by addition of the fluxes of all ionic species times their charge: i =− FA

F ∂Φ

∑ z J ( x) = −v( x)∑ z C + RT ∂x ∑ z C D j

j

j

j

2 j

j

j

Because of electroneutrality, the first summation term on the right-hand side of this equation vanishes, which yields dΦ = i

RT

∑

dx (z F C j D j ) A 2 j

2

(1.136)

This gives the potential drop dΦ associated with the current i through a space element of cross section A and thickness dx. By comparison with Ohm’s law and the definition of the resistivity, it is seen that the resistivity ρ of the solution is given by Equation 1.137. The ensuing resistance RΩ for a solution of cross section A and length ℓ is obtained in Equation 1.138. ρ=

1

∑( )

 z 2 F 2 C D /RT  j j  j 

RΩ =

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ρℓ A

(1.137)

(1.138)

69

Basic Concepts

Such an expression justifies a posteriori the previous discussion about ohmic drop effect in electrochemical cells (see Section III.A.3). Indeed, for a given resistance of the solution in Equation 1.138, the ohmic drop across the cell is expressed as (1.139)

Φ Ω = RΩi

Thus, the larger the concentration of dissociated ions, the smaller the ohmic drop. This in turn depends on the nature and affinity of the solvent/supporting electrolyte system, as discussed earlier and in Chapter 7. Let us now discuss more precisely the relationship between current and mass transfer at the electrode surface. From Equation 1.134, it is seen that the flux of electroactive species at the electrode surface (x = 0) is given, in the absence of migration (viz., in the presence of an excess of supporting electrolyte [106]), by Equation 1.140:  ∂C  J x =0 = − D    ∂x  x = 0

(1.140)

 ∂C  i = −nFAJ x = 0 = −nFAD    ∂x  x = 0

(1.141)

When there is no accumulation of material at the electrode surface, this flux corresponds to an exchange of n electrons per molecule, that is, to a current given in Equation 1.141.* Thus, for the electrochemical reaction in Equation 1.142, the reactant flux is associated with an identical, but with an opposite sign, product flux, that is, R + ne ⇌ P  ∂[ R ]   ∂[ P ]  i = nFADR  = −nFADP     ∂x  x =0  ∂x  x =0

(1.142)

Both fluxes must also correspond to the rate of electron exchange at the electrode surface. Since this latter has already been shown to be given by the Butler–Volmer rate law in Equation 1.109, the current must fulfill simultaneously the set of the following three equations: i = nFAk 0e − αnF [ E − E

0

] /RT

([ R ]

x =0

− [ P ]x =0 e nF ( E − E

0

)/ RT

)

(1.143)

 ∂[ R ]  i = nFADR    ∂x  x =0

(1.144)

 ∂[ P ]  i = −nFADP    ∂x  x =0

(1.145)

* Note that we use the polarographic convention generally used in organic electrochemistry due to historical reasons. In this convention, a reduction current is counted positive and an oxidation one negative. The IUPAC convention is opposite, so all analytical formulations given here correspond to −i when using the IUPAC rules.

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Comparison of Equations 1.143 and 1.144, for example, allows us to discuss more quantitatively the relationship between mass transfer rate and reversibility or irreversibility of an electron transfer step (see Table 1.4). Indeed, by equalizing both equations, one readily obtains the following for E = E 0: 1−

[ P ]x =0 DR (∂[ R ]/∂x ) x =0 = [ R ]x =0 k 0 [ R ]x = 0

(1.146)

From this simple equation, it is seen that when k0 is much larger than the flux at the electrode, the right-hand side tends to be zero, and thus [P]x = 0 ≈ [R]x = 0, which corresponds to the fulfillment of a Nernst equation at the electrode surface for E = E 0. Conversely, when k0 is much smaller than the demanded flux, one obtains [P]x = 0 > [R]x = 0 or [P]x = 0 < [R]x = 0, according to the flux sign. Thus, to obtain [P]x = 0 ≈ [R]x = 0, one must impose an electrode potential E, such as (1 − e nF [ E − E

0

] /RT

)e − αnF ( E − E

0

) /RT

=

DR (∂[ R ]/∂x ) x =0 k 0 [ R ]x =0

(1.147)

that is, such that n(E − E 0) ≪ 0, which is highly cathodic with respect to E0 for a reduction and highly anodic versus E0 for an oxidation. 3.

Microscopic Origin of the Diffusion Coefficient: Mass Transfer Rates and Diffusion Layer Thickness Before concluding this section devoted to Fick’s first law, we discuss the microscopic origin of the ubiquitous diffusion coefficient D (compare Equations 1.133 through 1.137, 1.140 through 1.142, and 1.144 through 1.147). In the preceding presentation, D was introduced as D = γRT, where γ is the proportionality constant between the velocity of the molecule and the forces acting on it, based on semimacroscopic considerations. In the absence of any applied force, a molecule is nevertheless moving because of thermal agitation. Yet owing to frequent collisions with the surrounding molecules, free motions occur only over atomic distances. The overall molecular movement then resembles a random succession of linear segments of approximately identical length ℓ and duration τ. Over a time period t = mτ, the average 2 2 square displacement ∆ is determined via statistical equations to be ∆ = mℓ 2 = (ℓ 2 /τ)t [107]. Let us now suppose that a force is applied to the particle arising, for example, from a concentration gradient along the x axis. Consider a plane perpendicular to the x axis at x = 0, as shown in Figure 1.18a. 2 During a time interval of t seconds, a random-walking molecule covers a mean distance L = (∆ )1/ 2 . Thus, those molecules from the negative axis region that may cross the plane at x = 0 during the time duration t originate from a volume element, adjacent to the plane, of thickness L. Let C(−L/2) be the average concentration of molecules within this element. Similarly, those that may also cross the plane but originate from the positive axis pertain to an identical volume element of thickness L in which the average concentration is C(+L/2). Since movements in a positive or negative direction x

–L (a)

0

J(x+dx)

J(x–dx)

+L

x–dx

x

x+dx

(b)

FIgURE 1.18 Symbolic representations of (a) the mass fluxes through one unit area surface, and (b) the entering and outgoing mass fluxes in an elementary volume of solution of unit area cross section (see text).

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Basic Concepts

are equally probable, one half of the molecules contained in each of these elements of solution cross the plane. Thus, a unit surface area of the plane at x = 0 is crossed by (1/2)LC(−L/2) molecules moving toward positive values of x and (1/2)LC(L/2) molecules moving toward negative values. The flux of molecules, that is, the balanced number of molecules that have crossed the unit surface area of the plane at x = 0 per unit of time and are lost (algebraically) from the layer on the negative axis side, is then J x=0 =

(1/ 2)LC (− L / 2) − (1/ 2)LC ( L / 2) t

Noting that [C(L/2) − C(−L/2)]/L is the concentration gradient at x = 0, which is (dC/dx)x = 0, this equation is rewritten as J x =0 = −

L2  dC  2t  dx  x =0

which, owing to the definition of L, yields finally J x =0 = −

ℓ 2  dC  2τ  dx  x =0

and thus D=

ℓ2 2τ

(1.148)

Comparison of this equation with Equation 1.132b or 1.134 provides a microscopic meaning to the diffusion coefficient D in Equation 1.148, where ℓ is the average free motion length and τ is the average time duration between two collisions. Besides its intrinsic value, such a result allows one to estimate how far a molecule has moved over an average period of time t. Indeed, from the expressions of L and D, one readily obtains L = (2 Dt )1/ 2

(1.149)

First, this equation provides a very convenient means to estimate, for example, the average distance covered by a molecule during a time interval t. In other words, a molecule generated at time zero at an electrode reaches a distance of approximately δ = (2Dt)1/2 after a time t has elapsed, provided its movements are controlled only by diffusion. For this reason, δ is usually called the diffusion layer thickness. This figure has an extreme importance in evaluating the nature of mass transfer at an electrode. Indeed, as discussed earlier, transport is governed by diffusion–migration only when δ < δconv 2 / 2 D beyond and by convection when δ > δconv. This limit corresponds to a maximum time tmax = δconv which convection dominates. For t < tmax, mass transfer processes at the electrode are then governed by diffusion (and possibly, at low excesses of supporting electrolyte, by migration). Conversely, when t > tmax, only an extremely small part of the molecular trajectories is controlled by diffusion, convection being the main means of mass transfer. As is elaborated in the following, this simple notion is extremely important in discriminating transient electrochemical techniques (t ≤ tmax) from steadystate techniques (t > tmax), that is, diffusion-controlled from convection-controlled mass transfer. The second important property of Equation 1.149 is that it provides an estimate of the rate, in terms of a characteristic time θ, associated with mass transfer. Indeed, this is the time θ needed for

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a molecule to reach the electrode, that is, to cover the space interval in which the molecular concentration differs from that in the bulk. In transient methods, this time is identical to that elapsed since the beginning of the experiment, provided that it is lower than tmax = δ2conv /2 D. For steady-state methods, the length to be covered is δconv, and thus from Equation 1.149, it follows that θ = δ2conv /2D. The rate of mass transfer can be defined as 1/θ, since it is obviously equivalent to a first-order process (see Chapter 10 for a demonstration of this point). Yet in light of the previous discussion, it is preferable to think in terms of a characteristic time θ associated with a given electrochemical method rather than in terms of mass transfer rate, although this intuitive latter notion was extremely worthwhile up to this point.* 4. Electrochemical Homogeneous Kinetics: Fick’s Second Law Fick’s second law is based upon mass conservation at any point in time and space. To formulate this point more precisely, let us consider that the spatial distribution of the concentration C of the species of interest depends on a single direction x.† Consider now a cylindrical space element of length 2 dx and cross-sectional area A, centered at x, and let us designate C(x,t) the average concentration at time t within this element, as shown in Figure 1.18b. During time interval dt, an algebraic flux J[x−dx,t] enters the volume element through its boundary at x−dx.‡ Similarly, an algebraic flux J[x + dx,t] leaves the element through. During the same time interval, k p dt moles of the species have been produced, via chemical reactions within the volume element, where k p and k c represent the kinetic rates of production and consumption of the species considered. Thus, the number of moles of the species varied by dN during the time duration dt: dN = J ( x − dx, t ) Adt − J ( x + dx, t ) Adt + k p dt − k c dt

(1.150)

Dividing the dN expression by the volume of element, 2A dx, gives the concentration variation, that is, ∂C J [ x − dx, t ] − J [ x + dx, t ] kp kc =− + − ∂t 2dx 2 Adx 2 Adx The first term on the right-hand side is the derivative of the flux vis-à-vis x, whereas k p /2 Adx and k c /2 Adx relate to the concentration variations from chemical origin, denoted (∂C/∂t)chem. ∂C ∂J  ∂C  = − + ∂t ∂x  ∂t chem

(1.151)

Then, Equation 1.151 simply relates the overall time dependence of the concentration to the algebraically additive mass transfer and chemical components. In this respect, it is important to point out that the chemical term is identical to that obtained under homogeneous conditions for the same chemical sequence and identical composition. Thus, Equation 1.151 constitutes a generalization of the usual kinetic rate laws to conditions in which concentrations are not homogeneous. Although * Note also that all the preceding discussions immediately transpose when comparing the half-lifetimes t1/2 associated with chemical reactions to θ. In our opinion, this is more satisfactory since it avoids the necessity of defining a rate constant associated with mass transfer processes, which are essentially of a physical, not a chemical, nature. † Note that, as for Fick’s first law, all the demonstrations easily adapt to a 3D space. Yet since for most electrochemical methods the linear space approximation, that is, depending only on the normal distance from the electrode, is adequate, we restrict ourselves to this simpler case in this presentation. ‡ Note that the terms enter and leave are used with respect to the volume shown in Figure 1.18 because fluxes are oriented with respect to the x axis. Thus, J > 0 through a surface corresponds to a flux oriented toward positive x values, and the converse is also true.

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developed here in the context of electrochemical techniques, it is valid in any kind of chemical situation in which concentrations are not uniform, as nearly any heterogeneous reactant or phasetransfer boundaries. Under electrochemical conditions, that is, near an electrode surface, two limiting formulations are obtained for Equation 1.151. For distances from the electrode larger than δconv, the solution is macroscopically homogeneous [108]. Thus, ∂J/∂x = 0, and Equation 1.151 simplifies to the usual kinetic rate law in Equation 1.152: x > δconv :

0 < x < δconv :

∂C  ∂C  = ∂t  ∂t chem

∂C ∂ 2C  ∂C  = D 2 +  ∂t ∂x  ∂t chem

(1.152)

(1.153)

Conversely, within the stagnant layer and in the presence of a sufficient excess of supporting electrolyte, the flux is given in Equation 1.134. Introduction of this latter figure in Equation 1.151 yields finally Equation 1.153, which is also commonly designated as Fick’s second law for diffusion reaction. In the derivation of Equation 1.152, we neglected the quantity of the species produced or consumed at the electrode. In practice, this is true for microscale electrochemical experiments but is necessarily inexact for macroscale or exhaustive electrolysis. To derive the general equation, let us consider the total equation for the bulk solution. The chemical contribution is given in Equation 1.152. On the other hand, a flux of the species at the convection layer limit δ conv must be considered. This flux corresponds to an algebraic number dNelec of moles of the species produced, given as follows:* − dN elec = J ( x = δconv , t ) Adt

(1.154)

where A is the electrode surface area. Dividing both members of Equation 1.154 by the volume V of the bulk solution allows the reformulation of Equation 1.154 as DA  ∂C   ∂C   ∂t  = − V  ∂x  −  elec   x =δconv

(1.155)

Thus, the overall concentration rate law in the solution is finally expressed as follows in a general case: x > δconv :

∂C DA  ∂C   ∂C  =− +    ∂t V  ∂x  x =δconv −  ∂t chem

(1.156)

The absolute value of the first term on the right-hand side of Equation 1.156 is at maximum equal to (DA/V)C0/δconv, where C0 is the bulk concentration of the electroactive reactant (see Section IV.B). The electrode consumption or production is then equivalent to pseudo-first-order kinetics whose apparent rate constant is at maximum equal to kelec =

DA Vδconv

(1.157)

− * The minus superscript in δconv means that the flux is that within the diffusion layer, that is, the limit of J(x), when x tends toward δ conv from lower values.

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To evaluate kelec, let us consider the following average figures: D = 10−5 cm2/s, δconv = 10 μm = 10−3 cm, and cell volume V = 100 cm3. Then, kelec = 10−4 A, when expressed in s−1 and A is expressed in cm2. For microscale electrochemical experiments, A is of the order of few square millimeters at most, that is, a few hundredths of a square centimeter; thus, kelec ≈ 10−6 s−1. This means that several hours of continuous electrolysis are needed to significantly perturb (>1%) the bulk solution. This is the reason why this term is neglected in practice, Equation 1.152 being used instead of the more rigorous Equation 1.156. Obviously, when preparative-scale electrolysis is considered, one wants to increase kelec in Equation 1.157, that is, use large A/V ratios and small δconv by increasing the electrode surface area and convection within the cell. Owing to the organization of this book, preparative-scale electrolysis is discussed separately in Chapter 10. Thus, in the following, we restrict to analytic conditions, viz., we assume that Equation 1.152 represents the dependence on time of the concentration. In other words, the electrochemical behavior, in the absence of migrational contributions, is given by a set of equations analogous to that in Equation 1.158, with boundary conditions at x = 0, depending on the electrode potential and kinetics (compare Equations 1.143 through 1.145), and at x = δconv, C = Cb, where the bulk concentration Cb is given as a function of time by Equation 1.159*: 0 < x < δconv :

∂C ∂ 2C  ∂C  = D 2 +  ∂t ∂x  ∂t chem

x > δconv : C = C b such as

dC b  dC b  =  dt  dt chem

(1.158)

(1.159)

5. Dimensionless Formulation of Electrochemical Equations The dimensionless analysis of equations is an analytic mathematical technique of frequent use in engineering. Indeed, this technique allows one to “concentrate” all the a priori independent parameters resulting in an identical effect into a single dimensionless parameter. To illustrate the principle and operational interest of the method, let us consider a general second-order homogeneous reaction, such as that in the following, the reactant concentration being C0 at the beginning (t = 0) of the experiment: 2R → P (k )

(1.160)

Strict application of homogeneous chemical kinetics readily yields the rate law in Equation 1.161, which governs the reactant concentration, as a function of time. Integration of this rate law gives the classic expression in Equation 1.162 for an experiment of duration θ. Yet Equation 1.162 may be rewritten under the equivalent form in Equation 1.163. d[ R ] = −2k[ R ]2 dt

(1.161)

1 1 = 0 + 2kθ [R] C

(1.162)

* Note that when a figure C depends on several independent variables x and t, its variations vis-à-vis one of these variables are given by its partial derivative ∂C/∂x with respect to the variable. When C depends on a single variable, the usual derivation is used: dC/dx. The intrinsic difference between the two kinds of derivations is indicated by the symbolic use of ∂ or d in the derivative notations.

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1 [R] = C 0 1 + 2kC 0θ

(1.163)

In many respects, the formulation in Equation 1.163 is more interesting than that in Equation 1.162. Indeed, its left-hand side gives the instantaneous fraction of the reactant not converted to the product at any time θ. Chemically, this is usually more significant than the actual concentration. Its righthand side shows that this fraction depends on a single factor (or parameter) 2kC0 θ, which includes the effect of the a priori three independent factors: an intrinsic factor k and two experimental parameters C0 and θ. Just at a glance it shows that doubling the concentration, for example, results in the same effect as keeping the same concentration and doubling the duration of the experiment, or increasing the temperature, so that k is doubled. In this example, the dimensionless reactant fraction r = [R]/C0 and dimensionless rate parameter λ = 2kC0 θ were introduced by reformulating the integrated rate law in Equation 1.162. However, this could have been done directly on the rate law in Equation 1.161 by introducing a dimensionless time τ = t/θ (i.e., the elapsed fraction of the overall experiment duration θ) and the reactant fraction r defined earlier. Then, Equation 1.161 reformulates as in Equation 1.164a and is associated with the initial condition in Equation 1.164b. Integration of Equation 1.164a, owing to the initial condition, yields readily Equation 1.165, which is identical to Equation 1.163. τ>0:

dr = − λ dτ r2

τ = 0: r =1

τ = 1: r =

1 1+ λ

(1.164a) (1.164b) (1.165)

The interest of this second approach, that is, dimensionless analysis of the rate law itself, becomes obvious for cases in which the integration of the rate law is not achievable analytically and must be performed numerically. Indeed, the set of Equations 1.164 shows that the variable of interest, r = [R]/C 0, depends only on λ = 2kC0 θ. Thus, would an analytic derivation be impossible to achieve, numerical integration must be performed only for a selected number of λ values to obtain a single curve giving r as a function of λ, that is, a function of all factors k, C0, and θ affecting the conversion yield. If one says that 10 different values of λ, for example, are needed to determine the curve with the required accuracy, the process needs 10 calculations. Yet without dimensionless analysis, the same information with the same accuracy would be contained in at least 1000 curves, that is, [R] as a function of 10 values of θ, for 10 selected values of C0 and for each 10 values of k. Besides the intrinsic interest of having all the information contained in a single curve, the method is essentially time saving. For example, consider that a numerical integration requires 30 s; the whole dimensionless curve needs about 5 min calculation time, whereas the normal or direct method requires at least 8.33 h! As explained in the previous parts of this section, two figures tightly control electrochemical kinetics: a spatial one δ, the thickness of the diffusion layer, and a temporal one θ, the time elapsed since the beginning of the electrochemical perturbation. Thus, in most situations, these figures constitute adequate and suitable references for the definitions of the dimensionless space y = x/δ and time τ = t/θ. Similarly, in most electrochemical kinetic experiments, the bulk concentration C0 of the starting material remains constant (see, however, Chapter 10 for preparative electrochemistry). It follows that s = [S]/C0 constitutes a meaningful dimensionless concentration for any species S whose real concentration is [S]. Another obvious dimensionless parameter concerns the potential figures. Indeed, in nearly all electrochemical circumstances, the electrode potential per se is not

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important but influences the electrochemical behavior through its difference from E 0 times nF/RT. (Compare, e.g., the Nernst equation or the Volmer–Butler rate law in Equation 1.143.) Thus, definition of the dimensionless potentials as ξ = (nF/RT)(E − E 0) appears suitable. y, τ, s, and ξ constitute the master dimensionless parameters from which the definition of all others ensues, as is shown in the following. Let us consider first the electrode boundary conditions in Equation 1.144 or 1.145. Introduction of r = [R]/C0 and y = x/δ in Equation 1.144, for example, directly affords nFADC 0  ∂r    δ  ∂y 0

i=

(1.166)

where the subscript 0 is used to indicate at y = 0, which shows that a convenient dimensionless definition for the current is ψ = i/(nFADC0/δ), which yields Equation 1.167:  ∂r  ψ=   ∂y 0

ψ=

(1.167)

k 0δ αξ e (r0 − p0e − ξ ) D

(1.168)

(1.169)

ψ = Λeαξ (r0 − p0e − ξ )

Introduction of r, p = [P]/C0, ξ = (nF/RT)(E − E0), and ψ in the Volmer–Butler rate law (Equation 1.143) readily yields Equation 1.168. The latter shows that a convenient dimensionless rate of electron transfer is Λ = k0 δ/D, since it compares the intrinsic value of the rate constant to that of the mass transfer process. Thus, Equation 1.168 reformulates as Equation 1.169. Let us now examine the time- and space-dependent partial derivative equation of the kind demonstrated by Equation 1.158, which describes variations in the concentration profiles in the stagnant layer adjacent to the electrode. For any species S, introducing τ, y, and s leads to reformulation of Equation 1.158 as in follows: 0 0) or anodic of E 0 for an oxidation process. The E1/2 represents both thermodynamic (E0) and kinetic contributions (lnk0 δ/D) and must not be used as a substitute for E 0. Figure 1.19a presents the variations in the steady-state voltammograms for the R/P redox system as a function of Λ = k0 δ/D for α constant and equal to 0.5. The variations in E1/2 with the same parameter k0 δ/D are presented in Figure 1.19b for selected constant values of α. Therefore, in an actual experiment, variations in E1/2 with δ (note that k0 and D are intrinsic figures for a given couple R/P and a given medium) are indicative of kinetic complications and should warn against the use of E1/2 as an estimate of E 0. 2. Mechanism Involving a Follow-Up Reaction: EC Mechanisms Let us first consider the case in which the initial electron transfer in Equation 1.190 is sufficiently rapid to be nernstian (i.e., Λ = k0 δ/D ≫ 1) but is followed by an irreversible pseudo-first-order chemical reaction (i.e., k may include the concentrations of other reagents, when constant). R + ne ⇌ P ( E 0 , k 0 , α) P →⋯

(k )

(1.190) (1.191)

The dimensionless equations governing the R and P concentration profiles are obtained from Table 1.5, as in Equations 1.192 and 1.193:  d 2r =0  2 dy   2 d p  p = λ  dy 2

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0 < y tanh λ1/2, the corrective term to E 0 in Equation 1.204 always has the sign of n. Thus, the wave is observed in a potential range E W such that n(E W − E 0) > 0, that is, anodic to E 0 for a reduction (n > 0) and cathodic to E 0 for an oxidation (n < 0). E1/ 2 = E 0 +

RT RT ln λ − ln tanh λ1/ 2 2nF nF

(1.204)

When λ → 0, that is, k ≪ D/δ2, one obtains E1/2 ≈ E 0, since then tanh λ1/2 ≈ λ1/2, resulting in a negligible corrective term in Equation 1.204. This means that the chemical reaction in Equation 1.191 is too slow vis-à-vis the mass transfer rate to significantly affect the voltammogram. The system is then said to be nernstian and chemically reversible. Conversely, when k ≫ D/δ2, that is, when the chemical reaction is faster than the mass transfer at the electrode, λ → ∞. Then, because tanh λ1/2 → 1 for λ1/2 → ∞, Equation 1.204 simplifies to the following: E1/ 2 = E 0 +

RT kδ2 ln 2nF D

(1.205)

From the latter, it is seen that for n = 1 and room temperature, for example, E1/2 shifts anodically by approximately 30 mV per unit of logkδ2/D. As for the slow electron transfer case, E1/2 is therefore not an acceptable substitute for E0 since it incorporates kinetic contributions that then arise from the follow-up chemical reaction. The variations in E1/2 with λ = kδ2/D result from the fact that when λ increases, the P concentration at the electrode surface decreases, as is apparent in Equation 1.201 and Figure 1.20.* Thus, at any driving force imposed by the value of ξ, one obtains from Equation 1.195 for large λ values: r0 =

ψ ( −ξ) ψ e = e( − ξ ) 1/ 2 2 λ (kδ /D)1/ 2

(1.206)

which shows that the reactant concentration is considerably less than that, r0 = ψe(−ξ), obtained at the same potential in the absence of a fast follow-up chemical step. Thus, the heterogeneous electron * Indeed from Equation 1.201, one obtains p 0 ≈ ψ/λ1/2 when λ > 7, with an accuracy better than 1%.

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1

1 .01 .1

[P]/C 0

[P0]/C 0

1

100 1000

10 0 0.5

0

1.0

0

4

log (kδ /D)

x/δ (a)

2 2

(b)

FIgURE 1.20 Steady-state electrochemical method. (a) Concentration profiles of the product obtained upon electron transfer in the EC sequence in Equations 1.190 and 1.191 as a function of the dimensionless chemical rate constant kδ2/D (numbers on the solid curves). The reactant concentration is shown for comparison as the dashed line. (b) Variations in the product electrode concentration as a function of kδ2/D. The dashed curve corresponds to the approximation in Equation 1.206.

transfer is continuously displaced by P removal in a way similar to that already observed for the corresponding homogeneous situation in Equations 1.207 and 1.208. Yet in the latter case, as elaborated in Section II.B (see Equations 1.45 and 1.47), the follow-up chemical step affects the overall kinetics by a factor involving k with a unity exponent, whereas an exponent of one-half is involved in the heterogeneous equivalent case: M + ne ⇌ N N + R ⇌ M + P (K )

(1.207)

P →⋯

(1.208)

(Compare the half-factor in Equation 1.205 or the half-exponent in Equation 1.206.) This effect, which arises from the heterogeneous nature of the electrochemical process (i.e., a surface reaction vis-à-vis a volume reaction in homogeneous phases), is the basis of the efficiency of redox catalysis or mediated electron transfer (see Section III.E.3). Thus, for a given redox system, as in the sequence in Equations 1.190 and 1.191, the use of a redox mediator M in Equation 1.207 allows the reduction of R to be performed at potentials less cathodic than E1/2 in Equation 1.205 (or the R oxidation at potentials less anodic than E1/2) for the same electrochemical setup (i.e., an identical mass transfer rate). 3.

Chemical Kinetics from Half-Wave Potentials: Determination of Rate Constants and Reaction Orders The preceding two examples were presented to illustrate the simplicity of electrochemical analysis under steady-state conditions. Indeed, the kinetic derivations are very similar to those encountered in the homogeneous chemical situation, except for the replacement of the usual derivative dC/dt by a second-order derivative vis-à-vis the space variable d2C/dt2 as apparent, for example, in Equation 1.193. The existence of a second-order differentiation introduces unusual dependencies on reaction orders when compared with those observed in homogeneous kinetics. Thus, if one considers the rate constant k in the chemical reaction (1.191) to be a pseudo-first-order rate constant, that is, k = k0[Z], where Z is a coreagent, the concentration of which is maintained constant, a dependency

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on [Z]1/2 is observed (because of the involvement of k with a power of one-half in Equation 1.198) for electrochemical experiments rather than the dependency on [Z] that would be observed under homogeneous experiments. Let us consider a general case in which the product formed upon electron transfer reacts via a chemical reaction of the ρth order in P and involving other coreagents A, B, …, in excess, as featured in the reaction sequences (1.209) and (1.210): R + ne ⇌ P (E 0 ) ρP + αA + βB + … → …

(1.209) (k0 )

(1.210)

One obtains the pseudo-rate constant k = k0[A]α[B]β…. Thus, dimensionless analysis affords the differential equation in Equation 1.211, which describes the concentration profile of the P species when λ is given as in Equation 1.212: d2 p = ρλpρ dy 2

λ = k (C 0 )(ρ−1)

δ2 , k = k0 [ A]α [ B]β …, D

(1.211)

(1.212)

When λ is extremely large and the usual situation in which the bulk solution contains no P is considered, integration of Equation 1.211 is performed as follows. Multiplication of both members by 2(dp/dy) dy yields 2

dp d 2 p dy = 2ρλpρdp dy dy 2

(1.213)

The left-hand term in Equation 1.213 is the derivation of (dp/dy)2 vis-à-vis y. The right-hand side is the derivation of [2ρλ/(ρ + 1)]p(ρ + 1) versus p. Thus, integration of Equation 1.213 yields 2

 dp  2ρλ ρ+1 p + cst   = dy ρ +1  

(1.214)

The integration constant in Equation 1.214 is determined by the condition p 0 ≈ 0 when y ≈ 1. When λ is large, this is also equivalent to saying that dp/dy ≈ 0 when y ≈ 1, since P exists only in a thin layer (called the kinetic layer) adjacent to the electrode (compare, e.g., Figure 1.20a). The integration constant in Equation 1.214 is then zero, and one obtains at the electrode surface, for y = 0, ψ2 =

2ρλ (ρ+1) p0 ρ +1

 dp  since ψ = −    dy 0

(1.215)

On the other hand, since R is not involved in chemical reactions, one obtains as in Equation 1.183, r0 = 1 − ψ. Thus, for the nernstian electron transfer in Equation 1.209, one obtains Equation 1.216 when taking Equation 1.215 into account. r0 = p0e( − ξ)

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 ρ +1  1 − ψ = ψ 2 /(ρ+1)    2λρ 

e( − ξ )

(1.216)

λ being a constant term, it may be included in the exponential factor, which yields the final equation of the steady-state voltammogram in Equation 1.217. From this equation, it is seen that the current depends only on the apparent dimensionless potential ξ* given in Equation 1.218. (1.217)

ψ + ψ 2 /(ρ +1)e( − ξ* ) = 1 ξ* = ξ +

1 2λρ nF 0 1 2λρ ln = (E − E ) + ln ρ +1 ρ +1 ρ + 1 ρ + 1 RT

(1.218)

From the equation ψ = 1 − r 0, the limiting plateau current corresponds to ψ = 1. Thus, the half-wave potential E1/2 is obtained by replacing ψ by 1/2 in Equation 1.217, which yields E1/ 2 = E 0 +

2ρ RT 1 − ρ RT /nF ln 2 + ln ρ +1 ρ +1 nF 1 + ρ

 RT 1   δ2  + ln k0 [ A]α [ B]β …(C 0 )(ρ −1)   D  nF (ρ + 1)  

(1.219)

Because of the expression of the last term, it is seen that at room temperature, for example, where RT/F ln10 ≈ 60 mV, E1/2 varies by [60α(ρ + 1)/n], [60β(ρ + 1)/n], and [60(ρ−1)(ρ + 1)/n] mV per unit of log[A], log[B], and logC0, respectively. These variations then allow an easy determination of the various reaction orders of any follow-up chemical reactions, provided the initial electron transfer is fast enough and the number n of electrons exchanged is known. Another approach consists of varying a reactant concentration and modifying the mass transfer rate δ in order that E1/2 remains constant. The term in braces in Equation 1.219 must then be kept constant; this is achieved when δ is equal to δ1/2, as follows: 1/ 2

k  δ1/ 2 = constant [ A]α / 2 [ B]β / 2 ⋯(C 0 )(ρ−1)/ 2  0  D

(1.220)

Expression (1.220) shows that the slope of logδ1/2 versus log[A] is α/2. Similarly, β/2 and (ρ − 1)/2 are the slopes observed for the variations with log[B] or logC0, which is a convenient way to determine the reaction orders. Note that although developed here for E1/2, this procedure may be used with any chosen potential E ε such that i/ilim has a fixed value ε. Similarly, variations in k0/D with the temperature (i.e., activation energy determinations) can be obtained through this procedure, provided care is taken of the RT/F factors, which need to be corrected (compare Equation 1.219) (see, e.g., Chapter 2 for extensions of this method). All these methods stem from the fact that for a given reaction sequence, such as that in Equations 1.209 and 1.210, which involve a single RDS, all the kinetics and thus the shape and location of the voltammogram depend only on the dimensionless rate constant parameter λ in Equation 1.212. As a result, any modification of the experimental conditions that keep λ constant does not modify the dimensionless voltammogram.* Thus, quantitative information on the chemical mechanism * Yet any variation(s) associated with the temperature may result in additional effects because of the temperature involvement in Equation 1.219.

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(Equation 1.210 may be a succession of chemical steps) is obtained without mathematical derivation, but only from dimensionless analysis (compare Chapter 2). However, determination of the intrinsic rate constant k0 or of the apparent ρth order rate constant k requires that the precise mathematical dependence of E1/2, for example, on the dimensionless rate constant be determined as in Equation 1.219. Indeed, the principle of all methods for determining k0 or k consists of increasing the mass transfer rate (i.e., here upon decreasing δ) up to a range such that kδ2/D ≈ 0, that is, E1/2 ≈ E 0. These methods constitute the frame on which any particular method can be elaborated. Yet in practice, the experimental difficulty is that with standard apparatus, δ2/D cannot be varied over an extremely wide range. For example, with the rotating disk electrode, which is the most convenient steady-state method (with the exception of ultramicroelectrodes [110]), δ depends on the rotation frequency ω of the electrode (see Chapter 2). Yet to maintain correct hydrodynamic conditions, ω cannot be varied, with standard apparatus, outside the range of 10 rotations per second (rps) to 1000 rps, which limits access to fast kinetics. Transient electrochemical techniques [102,111] allow this range of investigation to be widely extended (from 1 s to approximately 10 ns). Indeed, the same method may be used in a time domain extending over approximately eight orders of magnitude. Besides this very important point, transient electrochemical techniques provide current–time and current–potential patterns that are easily recognizable, for example, in cyclic voltammetry. Thus, with minimal eye training, a large amount of qualitative or nearly quantitative data may be obtained just by inspection of a transient voltammogram. This is in many respects an important advantage for diagnosis of kinetics, analogous with that of IR, ultraviolet, or nuclear magnetic resonance spectroscopies for structural information. In our opinion, this is one of the main reasons why transient electrochemical methods have progressively supplanted steady-state methods. Indeed, as illustrated by the preceding examples, steadystate voltammograms have generally sigmoidal shapes that are difficult to relate to a particular mechanism without quantitative analysis.

C.

TRANSIENT ELECTROCHEMICAL METHODS

1. Introduction: Time Hysteresis in Current Reversal Techniques As explained earlier, in transient electrochemical methods, an electrical perturbation (potential, current, charge, and so on) is imposed at the working electrode during a time period θ (usually less than 10 s) short enough for the diffusion layer δ ≈ (2Dθ)1/2 to be smaller than the convection layer δconv imposed by natural convection. Thus, the electrochemical response of the system investigated depends on the exact perturbation as well as on the elapsed time. This duality is apparent when one considers a double-pulse potentiostatic perturbation applied to the electrode as in the double-step chronoamperometric method. R + ne ⇌ P ( E 0 ) (n > 0)

(1.221)

Let us consider, for example, the simple nernstian reduction (n > 0) reaction in Equation 1.221 and a solution containing initially only the reactant R. Before any electrochemical perturbation, the electrode rest potential E1 is made largely positive to E 0. At time zero, the potential is stepped to a value E2, sufficiently negative to E 0, so that the concentration of R is close to zero at the electrode surface. After a time θ, the electrode potential is stepped back to E1, so that the concentration of P at the electrode surface becomes zero. When this potentiostatic perturbation, represented in Figure 1.21a, is applied in a steady-state method, the R and P concentration profiles are linear and depend only on the electrode potential but not on time, as shown in Figure 1.20a (for k ≈ 0). Yet when the same perturbation is applied in transient methods, the concentration profiles are curved and time dependent, as evidenced in Figure 1.21b. Thus, it is seen from this figure that a return

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0

1 1

5

E2 E0

0 1

∞ 20

0 C/C0

E1

15 11

E1

E1

Current

10 ms

C/C0

E2

0 0

θ

(a)

Time

0 (b)

20

0

x, μm

θ

Time

(c)

FIgURE 1.21 Transient potentiostatic electrochemical perturbation of a simple electron transfer reaction: (a) imposed potential; (b) resulting concentration profiles for the reactant (solid curves) or the product (dashed curves) for a step duration θ = 10 ms at various times from the beginning of the pulse; (c) resulting variations in the current.

step duration at E1 much longer than the step duration θ at E2 is needed for the initial concentration profiles to be restored. This hysteresis corresponds to the propagation of the diffusion perturbation within the solution, which then keeps a “memory” of the past perturbation [108b]. This information is stored via the structuring of the concentrations in the space near the electrode as a function of the elapsed time. The current flowing through the electrode is proportional to the gradient of the R (or the P) concentration profile at the electrode surface. Thus, the progressive smoothing of the concentration profiles with the time elapsed after each discontinuity of the perturbation results in a constant decrease in the current with time, as represented in Figure 1.21c. Let us now assume that the product P formed upon the electron transfer in Equation 1.221 may react chemically with a half-life t1/2. When θ ≪ t1/2, almost no P has time to disappear during the time duration of the experiment. The P concentration profile then remains identical to that shown in Figure 1.21b, and the current observed for t > θ, which corresponds to P reoxidation, is identical to that in Figure 1.21c. Conversely, when θ ≫ t1/2, nearly all P molecules produced at the electrode are consumed, and the P concentration profile is flat and close to zero, except in the close vicinity of the electrode. As a result, almost no oxidation current is observed for t > θ, that is, during the period in which the electrode potential is stepped back to E1 (see Figure 1.21a). In the intermediate range, that is, when θ and t1/2 are of similar orders of magnitude (generally 0.1 t1/2 < θ < 10 t1/2), intermediate values of the oxidation current are observed as shown in Figure 1.22. Determination of these variations allows the precise identification of the chemical follow-up sequence in which P is involved, as well as of the pertinent rate constants (see Chapter 2). However, this large sensitivity for mechanistic analysis is earned at the expense of two important factors. First, owing to the large potential variations, the fine dependence of the system on the electrode potential around E 0 is not seen. Second, and maybe more important, the current versus time variations have no important visual characteristics. Indeed, as seen from Figure 1.22, for example, all the curves are very similar, and only their juxtaposition on the same figure allows their shapes to be compared. For these two reasons, cyclic voltammetry has progressively supplanted potentiostatic or galvanostatic methods in electrochemical kinetic investigations. A more complete description of this method is given in Chapter 2, yet for our purposes, here we present the method briefly. It basically consists of applying a linear variation of the electrode potential E with time from a potential E i, where the reactant R is not electroactive, to a potential Ef sufficient for the reactant concentration to be zero at the electrode surface. At Ef, the potential is linearly scanned back to Ei, usually

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iθ1/2/(FSC 0D1/2) 2

1 ∞

0

10 1 0

–1 0

.5

1

2

1.5

Time, t/θ

FIgURE 1.22 Variations in the current trace in a potentiostatic experiment for an EC sequence for different values of the dimensionless chemical rate constant kθ (numbers on the curves).

with an identical slope with respect to the time. The latter, v, is designated as the scan rate (in volts per second). Figure 1.23a presents such a potential–time dependence, together with the corresponding i−E curve* (Figure 1.23b) obtained for the reduction process in Equation 1.221. Eye inspection of the cyclic voltammogram in Figure 1.23b shows the presence of two current peaks, in contrast to the monotonous variations of the current in Figure 1.21c. These peaks result from the existence of two opposite effects that successively control the magnitude of the current. For example, for the cathodic scan in Figure 1.23, the gradual decrease in R concentration at the electrode surface, when the potential is made more and more cathodic, tends to increase the current Ef Epc

E0

E0 v Ei 0 (a)

i R P

Time

Ef

Ei Ei

i Epa

Ei

Ei

E0

Ef

Ef

(c)

(b)

FIgURE 1.23 Cyclic voltammetry. (a) Imposed potential versus time variations. (b) Resulting transient current–potential curve for a simple electron transfer. The concentration profiles of the reactant R and product P are indicated at various characteristic potentials of the voltammogram. Epc and Epa, cathodic and anodic peak potentials. (c) Schematic change of the cyclic voltammogram as a function of the chemical stability of the product. * Note that in voltammetry (Figure 1.23), one presents usually the current as a function of the electrode potential rather than displaying it as a function of time as one does in chronoamperometry (Figure 1.22). Thus, in Figure 1.23b, the upper trace corresponds to the forward cathodic scan, whereas the lower trace corresponds to the backward anodic scan.

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by making the R concentration profile steeper. Yet the propagation of the perturbation into the solution, as shown in Figure 1.21b, tends to flatten the concentration profile, which results in a progressive decay of the current. In practice, it is easily understood that the large variations in the surface concentrations for potentials in the vicinity of E 0 overwhelm the diffusion effect, the current then tending to rise. Yet when [R]x = 0 is close to zero, the potential variations affect the current magnitude less than the diffusion propagation in the solution, which explains why the current progressively decays. For the backward scan, the same phenomenology applies to the product concentration [P]x = 0, which then results in the appearance of an inverted peak on the lower trace in Figure 1.23b. When P reacts chemically, the magnitude of the reverse peak gradually decreases (Figure 1.23c) because there is less and less P present in the solution. Again, this allows determination of mechanisms and their pertinent rate constants, as explained for the double-step chronoamperometric method. From this brief presentation, it is seen that an important aspect of cyclic voltammetry is that the shapes of cyclic voltammograms are extremely indicative of the chemical processes occurring at the electrode or in the solution. As such, it is an extremely useful tool for kinetic diagnosis. Yet the much more complicated shapes, when compared with those in Figure 1.22, for example, makes quantitative information on the current difficult to obtain. 2. Transient Electrochemical Methods and Chemical Kinetics From the preceding presentation, it is seen that the current variations with time or potential–time are intimately related to the concentration profiles of the species engaged in heterogeneous electron transfers at the electrode surface [110]. Thus, any kinetic or thermodynamic perturbation of these concentration profiles results in a variation of the current/time or current/potential–time transient characteristics as outlined in Figures 1.22 or 1.23c. As was also apparent in the preceding discussion, as well as in several other places in this chapter, the degree to which these perturbations affect the concentration profiles is a function of their relative effect vis-à-vis the mass transfer rate. Dimensionless analysis again proves to be a very convenient way to appreciate this degree of interference through the dimensionless rate constants λ, equilibrium constants κ, or heterogeneous rate constants Λ in Table 1.5. Thus, any experimental variation resulting in an overall constancy of these dimensionless parameters does not change the dimensionless current ψ versus τ or ψ versus ξ curve, that is, the dimensionless voltammogram. This is a situation identical to that developed more extensively for the steady-state methods and is the basis of most electrochemical methods for the determination of reaction orders and rate constants. Yet when applied to current reversal techniques, such as double-step chronoamperometry or cyclic voltammetry, these methods require that an appreciable current be observed during the backward perturbation, that is, for t > θ, in potentiostatic methods or after the potential scan inversion in cyclic voltammetry. This requires that the characteristic time θ of the method is adjusted to match the half-life t1/2 of the electrogenerated intermediate. Today, owing to the recent development of ultramicroelectrodes, θ can be routinely varied from a few seconds to a few nanoseconds [102]. Yet with basic standard electrochemical equipment, θ is usually restricted from the second to the low millisecond range. Thus, for experimental situations involving faster chemical reactions, current reversal techniques are of little use. Yet, as is the case for steady-state methods, much kinetic information allowing mechanistic discriminations and reaction order determinations may be gathered from characteristic potential changes. Obviously, staircase methods, such as chronoamperometries, which are blind to these variations, are not convenient, but such methods as cyclic voltammetry or its descendants (Chapter 2) are extremely precise and adequate through peak potential or half-peak potential analysis. Indeed, in cyclic voltammetry, peak potentials Ep play a role identical to that of half-wave potentials E1/2 in steady-state methods. As for the latter methods, peak potentials vary linearly with the logarithm of dimensionless kinetic parameters λ or Λ in Table 1.5, provided these latter have values

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sufficiently large when compared to unity [94]. These linear variations, which may be used for the determination of reaction orders, stem from the same mathematical reasons as explained in the case of E1/2. Yet, the physical reason is quite different as evidenced by the case of the simple EC sequence in the following: R + ne ⇌ P (E 0 ) P →⋯

(1.222) (1.223)

(k )

From Equation 1.174, the partial derivative equation describing the concentration profile of P is written in dimensionless terms (compare Table 1.5): 0< y:

∂p ∂ 2 p = − λp ∂τ ∂y 2

(1.224)

When λ = kθ ≫ 1, the kinetic term in Equation 1.224 tends to be extremely large. On the other hand, ∂p/∂τ cannot be infinite, for obvious physical reasons (except at possible discontinuities in the potential variations). Thus, the diffusion term ∂2p/∂τ2 must compensate for the kinetic term in order that ∂p/∂τ remain finite.* In other words, a quasi-steady state is reached by mutual compensation of kinetics and diffusion. As a result, ∂p/∂τ ≪ ∂2p/∂y2 and λp. Equation 1.224 thus becomes equivalent to Equation 1.225, at least as concerns the derivation of the concentration profile of P. 0< y: 0≈

∂2 p − λp ∂y 2

(1.225)

Comparison of Equation 1.225 and Equation 1.193 or 1.211 shows that the problem is identical to that presented for steady-state methods. Thus, the same mathematical derivations show that any characteristic potential figure (Ep or half-peak potential Ep/2, and so on) varies linearly with the logarithm of λ = kθ. Indeed, one obtains by simple transposition in Equation 1.219  RT  0 ( ρ−1) α β Ep = E 0 +   ln k0 [ A] [ B] ⋯(C ) θ + cst  nF (ρ + 1) 

{

}

(1.226)

for the peak potential variations relative to the EC sequence in the following: R + ne ⇌ P ( E 0 )

(1.227)

ρP + αA + βB + ⋯ → … (k0 ; k = k0 [ A]α [ B]β …)

(1.228)

In Equation 1.226, θ is usually defined as θ = RT/Fvn, v being the scan rate, yet any dimensionless parameter, such as θ = ΔE/v, where ΔE is an adequate difference of potential, is equally convenient. Then, even when k0 cannot be determined, its temperature variations, that is, ΔH#, as well as the different reaction orders ρ, α, β,…, may be determined through the same experimental procedures already discussed for the analogous case of E1/2 in steady-state methods. * Note that this may be established on firm mathematical grounds, but it is beyond the scope of this presentation [94].

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Extension of these approaches to much more sophisticated kinetic situations, which involve, for example, two or three routes for a key intermediate, has been developed (Reference 94 and references therein). Relative rate constants (or relations between rate constants) may then be obtained that allow, by reference to a known value, the absolute determination of all the series. For example, rate constants for the cleavage of aromatic halide anion radicals, corresponding to a half-life from about a second to a few picoseconds, have been determined using these procedures [103].

ACKNOWLEDgMENTS This work was supported by ENS, CNRS, and UPMC (UMR 8640 PASTEUR). The author wishes to acknowledge several discussions with Prof. Irina Svir and Dr. Alexander Oleinick and their useful comments and suggestions.

REFERENCES 1. Most of the concepts developed in this chapter are presented in detailed analytical form in: (a) Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Fundamentals & Applications, 2nd edn.; John Wiley & Sons: New York, 2001. (b) Kissinger, P. T.; Heineman, W. R., eds. Laboratory Techniques in Electroanalytical Chemistry; Marcel Dekker: New York, 1984. (c) Bockris, J. O’M.; Reddy, A. K. N. Modern Electrochemistry; Plenum: New York, 1977, Vols. 1, 2. 2. (a) Huffman, J. W. Acc. Chem. Res. 1983, 16, 399–405. (b) Rautenstrauch, V.; Geoffroy, M. J. Am. Chem. Soc. 1976, 98, 5035–5037. (c) Rautenstrauch, V.; Geoffroy, M. J. Am. Chem. Soc. 1977, 99, 6280–6286. (d) for a recent analogy in electrochemistry (termed “concerted electron transfer”) see, e.g.: Costentin, C.; Hajj, V.; Louault, C.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 2011, 133, 19160–19167. 3. See, e.g., (a) Shiner, V. J., Jr.; Whittaker, D. J. Am. Chem. Soc. 1963, 85, 2337–2338. (b) Warnhoff, E. W.; Reynolds-Warnhoff, P.; Wong, M. Y. H. J. Am. Chem. Soc. 1980, 102, 5956–5957. 4. See however, Ashby, E. C.; Goel, A. B.; Argyropoulos, J. N. Tetrahedron Lett. 1982, 23, 2273–2276 for the possible occurrence of paramagnetic intermediates (SET mechanisms [5]) in the reduction with metal alkoxides. 5. Ashby, E. C.; Boone, J. R. J. Am. Chem. Soc. 1976, 98, 5524–5531. 6. Wigfield, D. C.; Gowland, F. W. J. Org. Chem. 1977, 42, 1108–1109. 7. March, J. Advanced Organic Chemistry, 3rd edn.; John Wiley & Sons: New York, 1985, pp. 811–814. 8. For an extensive review on single electron transfer (SET) mechanisms, see Chanon, M; Tobe, M. L. Angew. Chem. 1982, 21, 1–23. 9. See Freeman, F.; Lin, D. K.; Moore, G. R. J. Org. Chem. 1982, 47, 56–59. 10. (a) Criegee, R.; Kraft, L.; Rank, B. Liebigs Ann. Chem. 1933, 507, 159–197. (b) Waters, W. A. Mechanisms of Oxidation of Organic Compounds; Wiley: New York, 1964, pp. 97–106. (c) See, however, Kochi, J. K. Organometallic Mechanisms and Catalysis; Academic Press: New York, 1978, p. 108, for the possible involvement of one-electron transfer steps, with, e.g., Vv/Viv. 11. Reference 10c, pp. 11–13. 12. Taube, H.; Myers, H. J. Am. Chem. Soc. 1954, 76, 2103–2111. 13. (a) Laviron, E. J. Electroanal. Chem. 1981, 124, 19–33. (b) Huang, Y.-F.; Wu, D.-Y.; Wang, A.; Ren, B.; Rondinini, S.; Tian, Z.-Q.; Amatore, C. J. Am. Chem. Soc. 2010, 132, 17199–17210. (c) Utley, J. H. P. In: Weinberg, N. L. (ed.); Technique of Electroorganic Synthesis; Wiley Interscience: New York, 1974, Part 1, Chapter 6. 14. See e.g., Reference 2d and: Andrieux, C. P.; Savéant, J.-M.; Su, K. B. J. Phys. Chem. 1986, 90, 3815– 3823 and references therein as well as Chapter 24, this book. 15. (a) Savéant, J.-M. J. Am. Chem. Soc. 1987, 109, 6788–6795. (b) Wentworth, W. E.; George, R.; Keith, H. J. Chem. Phys. 1969, 51, 1791–1801. (c) Symons, M. C. R. Pure Appl. Chem. 1981, 53, 223–238. 16. Andrieux, C. P.; Savéant, J.-M.; Zann, D. Nouv. J. Chim. 1984, 8, 107–116. 17. See e.g., Kochi, J. K. Reference 10c. 18. See Chapters 15 and 43 in this book. 19. See Reference 1c, pp. 1124–1129. 20. See Reference 1a, pp. 44–63, for a thermodynamic derivation of Nernst equation.

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21. (a) See Reference 1a, pp. 53–54. (b) Ives, D. J. G.; Janz, G. J., eds. Reference Electrodes; Academic Press: New York, 1961. (c) Charlot, G.; Badoz-Lambling, J.; Tremillon, B. Les Reactions Electrochimiques; Masson: Paris, France, 1959, pp. 147–149, 347–348. 22. (a) Reference 1a, pp. 63–74. (b) Reference 1c, Vol. 1, Chapter 4. (c) Lingane, J. J., Electroanalytical Chemistry, 2nd edn.; Wiley-Interscience: New York, 1958, Chapter 3. (d) Vetter, K. J., Electrochemical Kinetics; Academic Press: New York, 1967. 23. As, e.g., the Henderson equation, established in Reference 1a, p. 72. 24. See, e.g., Reference 1c, pp. 688–717, and references therein. 25. Peover, M. E. In: Bard, A. J. (ed.); Electroanalytical Chemistry; Marcel Dekker: New York, 1967, Vol. 2, pp. 1–51. 26. See Reference 1c, pp. 660–679. 27. For example, consider the evaluation of solvation energies via the Born model: (a) Reference 1c, pp. 49–60. (b) Latimer, W. M.; Pitzer, K. S.; Slansky, C. M. J. Chem. Phys. 1939, 7, 108–111. (c) Noyes, R. M. J. Am. Chem. Soc. 1962, 84, 513–522. (d) Coetzee, J. F.; Campion, J. J. J. Am. Chem. Soc. 1967, 89, 2513–2517. (e) Tanaka, M. Inorg. Chem. 1976, 15, 2325–2327. 28. Dewar, M. J. S.; Hashmall, J. A.; Trinajstic, N. J. Am. Chem. Soc. 1970, 92, 5555–5559. 29. Howell, J. O.; Goncalves, J. M.; Amatore, C.; Klasinc, L.; Wightman, R. M.; Kochi, J. K.; J. Am. Chem. Soc. 1984, 106, 3968–3976. 30. For an extensive series of E 0 values, see Meites, L.; Zuman, P. CRC Handbook Series in Organic Electrochemistry; CRC Press: Cleveland, OH, 1975, Vols. I, II; and Meites, L.; Zuman, P.; Rupp, E. B. CRC Handbook Series in Organic Electrochemistry; CRC Press: Cleveland, OH, 1975, Vols. III–V. 31. (a) Steitwieser, A. Jr. Molecular Orbital Theory for Organic Chemists; Wiley: New York, 1961. (b) Zuman, P. Substituent Effects in Organic Polarography; Plenum: New York, 1967. (c) Hedges, R. M.; Matsen, F. A. J. Chem. Phys. 1958, 28, 950–953. (d) Hoijtink, G. J. Rec. Trav. Chim. Pays-Bas 1955, 74, 1525–1539. (e) Hoijtink, G. J. Rec. Trav. Chim. Pays-Bas 1958, 77, 555–558. (f) Schmidt, R. M.; Heilbronner, E. Helv. Chim. Acta 1954, 37, 1453–1466. (g) Parker, V. D. J. Am. Chem. Soc. 1976, 98, 98–103. 32. (a) Zuman, P. In: Progress in Physical Organic Chemistry; Streitwieser, A., Jr.; Taft, R. W. (eds.); WileyInterscience: New York, 1967, Vol. 5. (b) Reference 31b, p. 72. 33. Reference 1a, pp. 52–53. 34. For the dependence of activity coefficients on experimental conditions, see, e.g., (a) Reference 1c, pp. 180–267. (b) Justice, M. C.; Justice, J. C.; J. Sol. Chem. 1976, 5, 543–561. 35. See, e.g., Mansfield, W. C. Oxidation-Reduction Potentials of Organic Systems; Williams & Wilkins: Baltimore, MD, 1960, pp. 118–145. 36. (a) Schlesener, C. J.; Amatore, C.; Kochi, J. K. J. Am. Chem. Soc. 1984, 106, 3567–3577. (b) Schlesener, C. J.; Amatore, C.; Kochi, J. K. J. Am. Chem. Soc. 1984, 106, 7472–7482. (c) Schlesener, C. J.; Amatore, C.; Kochi, J. K. J. Phys. Chem. 1986, 90, 3747–3756. 37. (a) Wong, C. L.; Kochi, J. K. J. Am. Chem. Soc. 1979, 101, 5593–5603. (b) Reference 36a. 38. Compare (a) Parker, V. D. J. Electroanal. Chem. 1969, 22, A1–A3. (b) Bewick, A.; Mellor, J. M.; Pons, B. S. Electrochim. Acta 1978, 23, 77–79. 39. (a) Reference 38b. (b) Rollick, K. L.; Kochi, J. K. J. Am. Chem. Soc. 1982, 104, 1319–1330. 40. Eberson, L.; Jönsson, L.; Wistrand, L. G. Acta Chem. Scand. 1978, B32, 520–530. 41. Marcus, R. A. J. Chem. Phys. 1956, 24, 966–978. For quadratic relations arising from different origins, see also, e.g., (a) Magnoli, D. E.; Murdoch, J. R. J. Am. Chem. Soc. 1981, 103, 7465–7469. (b) Murdoch, J. R.; Magnoli, D. E. J. Am. Chem. Soc. 1982, 104, 3792–3800. (c) Murdoch, J. R. J. Am. Chem. Soc. 1983, 105, 2159–2164. (d) Murdoch, J. R. J. Am. Chem. Soc. 1983, 105, 2667–2672; (e) Kurz, J. G. Chem. Phys. Lett. 1978, 57, 243–246. (f) Grunwald, E. J. J. Am. Chem. Soc. 1985, 107, 125–133. 42. For a specialized and extensive review, see, e.g., Sutin, N. In: Lippard, S. J. (ed.); Progress in Inorganic Chemistry; Wiley-Interscience: New York, 1983, Vol. 30, pp. 441–498. 43. (a) Marcus, R. A. Annu. Rev. Phys. Chem. 1964, 15, 155–196. (b) Marcus, R. A. J. Chem. Phys. 1965, 43, 679–701. (c) Waisman, E.; Worry, G.; Marcus, R. A. J. Electroanal. Chem. 1977, 82, 9–28. (d) Marcus, R. A. Faraday Discuss. Chem. Soc. 1982, 74, 7–15. (e) Marcus, R. A.; Sutin, N. Biochim. Biophys. Acta 1985, 811, 265–322. 44. (a) Levich, V. G. In: Delahay, P. (ed.); Advances in Electrochemistry and Electrochemical Engineering; Wiley-Interscience: New York, 1966, Chapter 4. (b) Vorotyntsev, M. A.; Dogonadze, R. R.; Kuznetsov, A. M. Dokl. Akad. Nauk SSSR 1970, 195, 1135–1138. (c) German, E. D.; Dvali, V. G.; Dogonadze, R. R.; Kuznetsov, A. M. Elektrokhimiya 1976, 12, 667–672; Sov. Electrochem. 1976, 12, 639–643.

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45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55.

56. 57. 58. 59. 60. 61. 62. 63. 64. 65. 66. 67. 68. 69. 70. 71. 72. 73. 74. 75. 76.

Organic Electrochemistry (d) Dogonadze, R. R.; Kuznetsov, A. M. Prog. Surf. Sci. 1975, 6, 1–41. (e) Dogonadze, R. R.; Kuznetsov, A. M.; Levich, V. G. Electrochim. Acta 1968, 13, 1025–1044. (f) Dogonadze, R. R. In: Hush, N. S. (ed.); Reactions of Molecules at Electrodes; Wiley-Interscience: New York, 1971, Chapter 3. (a) Debye, P. Trans. Electrochem. Soc. 1942, 82, 265–272. (b) v. Smoluchowski, M. Z. physik. Chem. 1917/1918, 92, 129–168; volume 92 was dated 1918 and contains issues from 1916 to 1918, the referred paper appeared in an issue printed in 1917. Marcus, R. A. Discuss. Faraday Soc. 1960, 29, 21–31. Compare, e.g., Reference 36c. Libby, W. F. J. Phys. Chem. 1952, 56, 863–868. Atkins, P. W. Physical Chemistry, 7th edn; Oxford University Press: Oxford, U.K., 2002, pp. 956–961. Reference 49: (a) p. 958; (b) pp. 820–824. See, e.g., (a) Brown, G. M.; Sutin, N. J. Am. Chem. Soc. 1979, 101, 883–892. (b) North, A. M. The Collision Theory of Chemical Reactions in Liquids; Wiley: New York, 1964. See, e.g., References 44a and 44f, or Sutin, N.; Brunschwig, B. S. ACS Sym. Ser. 1982, 198, 105–135 for evaluation of κ, as a function of Vi, on the basis of a Landau-Zener [53] treatment of the electron tunneling probability Pe. (a) Landau, L. Phys. Z. Sowjetunion 1932, 2, 46–51. (b) Zener, C. Proc. R. Soc. Lon. Ser-A 1932, 137, 696–702. (c) Zener, C. Proc. R. Soc. Lon. Ser-A 1933, 140, 660–668. See, e.g., Reference 15a for a simple approach to the problem when k = 1. For more elaborate models, see, e.g., (a) Reference 44b. (b) Cannon, R. D. Chem. Phys. Lett. 1977, 49, 299–304. (c) Cannon, R. D. Electron Transfer Reactions; Butterworths: London, U.K., 1980. (d) Peover, M. E.; Powell, J. S. J. Electroanal. Chem. 1969, 20, 427–433. (e) Falsig, M.; Lund, H.; Nadjo, L.; Savéant, J.-M. Nouv. J. Chim. 1980, 4, 445–452. (f) As well as a critical review in German, E. D.; Kuznetsov, A. M. Electrochim. Acta 1981, 26, 1595–1608. See Reference 42, pp. 486–487. (a) Marcus, R. A. in Reference 41. (b) Marcus, R. A. J. Chem. Phys. 1957, 26, 867–871. (c) Marcus, R. A. J. Chem. Phys. 1957, 26, 872–877; (d) Hush, N. S. Trans. Faraday Soc. 1961, 57, 557–580; (e) Hush, N. S. J. Electroanal. Chem. 1999, 460, 5–29. See, e.g., Reference 42, pp 455–459, for a simple presentation of the problem on quantum mechanical grounds. See, e.g., Reference 1c, pp. 59–60 and pp. 201–202. (a) Efrima, S.; Bixon, M. Chem. Phys. Lett. 1974, 25, 34–37. (b) Efrima, S.; Bixon, M. J. Chem. Phys. 1976, 64, 3639–3647. (c) Efrima, S.; Bixon, M. Chem. Phys. 1976, 13, 447–460. (a) Marcus, R. A. J. Phys. Chem. 1968, 72, 891–899. (b) Agmon, N.; Levine, R. D. Chem. Phys. Lett. 1977, 52, 197–201. (c) Levine, R. D. J. Phys. Chem. 1979, 83, 159–170. See, e.g., Reference 42, pp. 479–481, for a more complete discussion. See, e.g., Eberson, L. Adv. Phys. Org. Chem. 1982, 18, 79–185. Rehm, D.; Weller, A. Isr. J. Chem. 1970, 8, 259–271. See, e.g., Meyer, T. J. In: Lippard, S. J. (ed.); Progress in Inorganic Chemistry; Wiley-Interscience: New York, 1983, Vol. 30, pp. 420–424 for a detailed discussion and pertinent references on the “inverted region.” Brunschwig, B. S.; Creutz, C.; Macartney, D. H.; Sham, T. K.; Sutin, N. Faraday Discuss. Chem. Soc. 1982, 74, 113–127. (a) Nakajima, T.; Toyota, A.; Kataoka, M. J. Am. Chem. Soc. 1982, 104, 5610–5616; (b) Iwasaki, M.; Toriyama, K.; Nunome, K. J. Chem. Soc. Chem. Comm. 1983, 320–322. (c) Salem, L. The Molecular Orbital Theory in Conjugated Systems; Benjamin: New York, 1966, pp. 467–485. Dickens, J. E.; Basolo, F.; Neumann, H. M. J. Am. Chem. Soc. 1957, 79, 1286–1290. See, e.g., Reference 1a, pp. 640–644, or Reference 1b, pp. 163–192, for descriptions of related electronics. (a) Howell, J. O.; Kuhr, W. G.; Ensman, R. E.; Wightman, R. M. J. Electroanal. Chem. 1986, 209, 77–90. (b) Amatore, C.; Jutand, A.; Pflüger, F. J. Electroanal. Chem. 1987, 218, 361–365. See Chapter 7, this book. Reference 1a, pp. 27–28. Amatore, C. In: Rubinstein, I. (ed.); Physical Electrochemistry: Principles, Methods, and Applications; Marcel Dekker: New York, 1995, Chapter 4. (a) Von Helmholtz, H. L. F. Ann. Phys.-Leipzig 1853, Ser. 2, 89, 211–233, continued on pp. 353–377. (b) Von Helmholtz, H. L. F. Ann. Phys.-Leipzig 1879, Ser. 3, 7, 337–382. (a) Gouy, G. J. Phys. Théor. Appl. 1910, 9, 457–468. (b) Chapman, D. L. Philos. Mag. 1913, 25, 475–481. See, e.g., Reference 1a, pp. 544–554, as well as references quoted therein, for a detailed presentation of the various models.

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77. Hale, J. M. In: Hush, N. S. (ed.); Reactions of Molecules at Electrodes; Wiley-Interscience: New York, 1971, Chapter 4, p. 229. 78. (a) Hush, N. S. J. Chem. Phys. 1958, 28, 962–972. (b) Reference 57c. (c) Hush, N. S. Electrochim. Acta 1968, 13, 1005–1023. (d) Hush, N. S. J. Electroanal. Chem. 1999, 470, 170–195. 79. (a) Thirsk, H. R.; Schmidt, P. P. In: Thirsk, H. R. (Senior Reporter); Electrochemistry: A Specialist Periodical Report; The Royal Society of Chemical: London, U.K., 1977, Vol. 5. (b) Thirsk, H. R.; Schmidt, P. P. In: Thirsk, H. R., Senior Reporter;  Electrochemistry: A Specialist Periodical Report; The Royal Society of Chemical: London, U.K., 1978, Vol. 6. 80. Kojima, H.; Bard, A. J. J. Am. Chem. Soc. 1975, 97, 6317–6324. 81. Yet the problem is not so clear-cut. See, e.g., Andrieux, C. P.; Blocman, C.; Dumas-Bouchiat, J. M.; Savéant, J.-M. J. Am. Chem. Soc. 1979, 101, 3431–3441. 82. (a) Butler, J. A. V. Trans. Faraday Soc. 1924, 19, 729–733. (b) Erdey-Gruz, T.; Volmer, M. Z. physik. Chem. (Leipzig) 1930, 150A, 203–213. 83. (a) Tafel, J. Z. physik. Chem. (Leipzig) 1905, 50, 641–712. (b) Reference 1a, pp. 103–105. 84. (a) Savéant, J.-M.; Tessier, D. J. Electroanal. Chem. 1975, 65, 57–66; (b) Savéant, J.-M.; Tessier, D. J. Phys. Chem. 1977, 81, 2192–2197. (c) Savéant, J.-M.; Tessier, D. J. Phys. Chem. 1978, 82, 1723–1727. (d) Savéant, J.-M.; Tessier, D. Faraday Discuss. Chem. Soc. 1982, 74, 57–72. (e) Garreau, D.; Savéant, J.-M.; Tessier, D. J. Phys. Chem. 1979, 83, 3003–3007. 85. (a) Nadjo, L.; Savéant, J.-M. J. Electroanal. Chem. 1973, 48, 113–145. (b) Klingler, R. J.; Kochi, J. K. J. Am. Chem. Soc. 1982, 104, 4186–4196. 86. See Reference 1a, pp. 554–569, and references therein, for a detailed presentation of adsorption thermodynamics and kinetics. 87. (a) Amatore, C.; Savéant, J.-M.; Tessier, D. J. Electroanal. Chem. 1983, 146, 37–45. (b) Amatore, C.; Savéant, J.-M.; Tessier, D. J. Electroanal. Chem. 1983, 147, 39–51, and references therein. 88. See, e.g., Deakin, M. R.; Stutts, K. J.; Wightman, R. M. J. Electroanal. Chem. 1985, 182, 113–122. 89. See, e.g., Ahlberg, E.; Parker, V. D. Acta Chem. Scand. 1983, B37, 723–730. 90. (a) Klymenko, O.V.; Svir, I.; Amatore, C. J. Electroanal. Chem., 2013, 688, 320–327. (b) Klymenko, O.V.; Svir, I.; Amatore, C. Mol. Phys., 2014, 112, 1273–1283. 91. Mann, C. K. In: Bard, A. J. (ed.); Electroanalytical Chemistry, Vol. 3; Marcel Dekker: New York, 1969, pp. 57–134. 92. (a) Geng, L.; Ewing, A. G.; Jernigan, J. C.; Murray, R. W. Anal. Chem. 1986, 58, 852–860. (b) Lines, R.; Parker, V. D. Acta Chem. Scand. Ser. B 1977, B31, 369–374. (c) Howell, J. O.; Wightman, R. M. Anal. Chem. 1984, 56, 524. (d) Howell, J. O.; Wightman, R. M. J. Phys. Chem. 1984, 88, 3915–3918. 93. See, e.g., (a) M’Halla, F.; Pinson, J.; Savéant, J.-M. J. Am. Chem. Soc. 1980, 102, 4120–4127. (b) M’Halla, F.; Pinson, J.; Savéant, J.-M. J. Electroanal. Chem. 1978, 89, 347–361. (c) Amatore, C.; M’Halla, F.; Savéant, J.-M. J. Electroanal. Chem. 1981, 123, 219–229. 94. (a) Savéant, J.-M.; Vianello, E. Electrochim. Acta 1963, 8, 905–923. (b) Savéant, J.-M.; Andrieux, C. P.; Nadjo, L. J. Electroanal. Chem. 1973, 41, 137–141. (c) Fatouros, N.; Chemla, M.; Amatore, C.; Savéant, J.-M. J. Electroanal. Chem. 1984, 172, 67–81. (d) Amatore, C.; Garreau, D.; Hammi, M.; Pinson, J.; Savéant, J.-M. J. Electroanal. Chem. 1985, 184, 1–24; (e) Andrieux, C. P.; Savéant, J.-M. In: Bernasconi, C. F. (ed.); Investigations of Rates and Mechanisms of Reactions, Vol. 6; John Wiley & Sons: New York, 1986, 41E, Part 2, Chapter 7. 95. (a) Andrieux, C. P.; Savéant, J.-M.; Su, K. B. J. Phys. Chem. 1986, 90, 3815–3823. (b) Reference 15a, and references therein. 96. (a) Amatore, C.; Savéant, J.-M. J. Electroanal. Chem. 1978, 86, 227–232. (b) Amatore, C.; Savéant, J.-M. J. Electroanal. Chem. 1979, 102, 21–40. 97. Smith, W.; Bard, A. J. J. Am. Chem. Soc. 1975, 97, 5203–5210. 98. For a review, see, e.g., Savéant, J.-M. Acc. Chem. Res. 1980, 13, 323–329, as well as Chapters 15, 43, and 44 in this book. 99. (a) Bunnett, J. F. Acc. Chem. Res. 1978, 11, 413–420. (b) Rossi, R. A.; Rossi, R. H. Aromatic. Nucleophilic Substitutions by the SRN1 Mechanism; ACS Monographs, 178, ACS: Washington, DC, 1983. 100. See, e.g., Alam, N.; Amatore, C.; Combellas, C.; Pinson, J.; Savéant, J.-M.; Thiebault, A.; Verpeaux, J. N. J. Org. Chem. 1988, 53, 1496–1504 and references therein. 101. See, e.g., [13a,b] and: (a) Laviron, E. J. Electroanal. Chem. 1983, 148, 1–16. (b) Laviron, E. J. Electroanal. Chem. 1984, 169, 29–46. (c) Laviron, E. J. Electroanal. Chem. 1986, 208, 357–372. (d) Lerke, S. A.; Evans, D. H.; Feldberg, S. W. J. Electroanal. Chem. 1990, 296, 299–315. (e) Hong, S. H.; Evans, D. H.; Nelsen, S. F.; Ismagilov R. F. J. Electroanal. Chem. 2000, 486, 75–84. (f) Macias-Ruvalcaba, N. A.; Evans, D. H. J. Phys. Chem. B 2006, 110, 24786–24795.

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102. Amatore, C.; Maisonhaute, E.; Simonneau, G. J. Electroanal. Chem. 2000, 486, 141–155. 103. Amatore, C.; Oturan, M. A.; Pinson, J.; Savéant, J.-M.; Thiebault, A. J. Am. Chem. Soc. 1985, 107, 3451–3459. 104. See, e.g., (a) Newman, J. S. Electrochemical Systems; Prentice Hall: Englewood Cliffs, NJ, 1973, pp. 305–339. (b) Levich, V. G. Physicochemical Hydrodynamics; Prentice Hall: Englewood Cliffs, NJ, 1962, and references therein. 105. Amatore, C.; Szunerits, S.; Thouin, L.; Warkocz, J.-S. J. Electroanal. Chem. 2001, 500, 62–70. 106. Amatore, C.; Deakin, M.R.; Wightman, R. M. J. Electroanal. Chem. 1987, 225, 49–63. 107. See, e.g., Reference 1a, pp. 128–130. 108. (a) Amatore, C.; Pebay, C.; Sella, C.; Thouin, L. ChemPhysChem 2012, 13, 1562–1568. (b) Amatore, C.; Klymenko, O. V.; Svir, I. Anal. Chem. 2012, 84, 2792–2798. 109. Reference 1c, p. 1069. 110. See, e.g., (a) Fleischmann, M.; Lasserre, F.; Robinson, J.; Swan, D. J. Electroanal. Chem. 1984, 177, 97–114. (b) Fleischmann, M.; Lasserre, F.; Robinson, J. J. Electroanal. Chem. 1984, 177, 115–127. (c) Amatore C., In: I. Rubinstein (ed.), Physical Electrochemistry: Principles, Methods and Applications; Marcel Dekker: New York, 1995, Chapter 4. 111. See Johnson, D. C.; Ryan, M. D.; Wilson, G. S. Anal. Chem. 1986, 58, 33R–49R, as well as Chapter 2 in this book.

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2

Techniques for Studies of Electrochemical Reactions in Solution Ole Hammerich and Bernd Speiser

CONTENTS I. II.

Introduction ............................................................................................................................ 98 Linear Sweep Voltammetry and Cyclic Voltammetry ........................................................... 98 A. Introduction ..................................................................................................................... 98 B. Experimental Setup....................................................................................................... 100 1. The Electrochemical Cell ...................................................................................... 100 2. The Solvent-Supporting Electrolyte System .......................................................... 101 3. The Electronic Equipment ..................................................................................... 102 C. Simple Electron Transfer Reactions.............................................................................. 102 1. Reversible Electron Transfer .................................................................................. 103 2. Quasireversible Electron Transfer.......................................................................... 106 3. Irreversible Electron Transfer ................................................................................ 107 4. Concluding Remarks and Examples ...................................................................... 108 D. Electron Transfer Reactions Followed by Chemical Reactions in Solution ................. 110 1. Typical Irreversible Follow-Up Reactions ............................................................. 110 2. Kinetic Classification of Simple Irreversible Follow-Up Reactions ...................... 111 3. Irreversible Follow-Up Reactions, Mixed Diffusion, and Kinetic Control (CV and DCV) ....................................................................................................... 113 4. Irreversible Follow-Up Reactions, Purely Kinetic Control (LSV) ........................ 117 5. Reversible Follow-Up Reactions ............................................................................ 119 6. Irreversible Follow-Up Reactions, the Prepeak Method ........................................ 124 7. Irreversible Follow-Up Reactions, Redox Catalysis .............................................. 125 8. Chemical Reactions Following Quasireversible or Irreversible Electron Transfer Reactions ................................................................................... 128 9. General Remarks and Conclusions ........................................................................ 128 E. Limiting Experimental Factors ..................................................................................... 128 1. The Cell and the Working Electrode ..................................................................... 129 2. Uncompensated Solution Resistance, Ru................................................................ 130 3. Importance of Precision in Potential Measurements ............................................. 132 4. Background Currents ............................................................................................. 132 F. Computer-Based Methods for Analysis of Voltammetric Data .................................... 132 1. Semi-Integration and Convolution Techniques ...................................................... 133 2. Fitting Simulated to Experimental Voltammograms ............................................. 133

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III.

Ultramicroelectrodes and Scanning Electrochemical Microscopy ...................................... 134 A. Ultramicroelectrodes .................................................................................................... 134 1. Determination of Eo′ for the A/B Redox Couple When B Is Highly Reactive...... 136 2. Determination of Heterogeneous Electron Transfer Rate Constant ...................... 138 3. Kinetics of Rapid Follow-Up Reactions ................................................................ 138 4. Voltammetry in Electrolytes with Low Conductivity............................................ 139 B. Application of Scanning Electrochemical Microscopy for Studies of Reaction Kinetics...................................................................................................... 139 IV. Potential-Step and Current-Step Methods ............................................................................ 141 A. Chronoamperometry and Double Potential-Step Chronoamperometry ....................... 141 B. Chronocoulometry and Double Potential-Step Chronocoulometry ............................. 144 C. Chronopotentiometry and Current-Reversal Chronopotentiometry ............................. 144 V. Polarography ......................................................................................................................... 146 VI. Methods Based on Forced Convection ................................................................................. 150 A. General Considerations .................................................................................................150 B. The Rotating Disk Electrode.........................................................................................150 C. The Rotating Ring-Disk Electrode ............................................................................... 154 VII. Methods for Determination of the Number of Electrons Transferred per Molecule of Substrate ........................................................................................................... 157 VIII. Methods for Determining the Diffusion Coefficient of Electroactive Molecules ................ 159 References ...................................................................................................................................... 160

I. INTRODUCTION Electrochemical methods are widely used to gain information about the kinetics and mechanisms of chemical reactions associated with the electron transfer at an electrode. A unique feature of these methods is that the electrode serves both as the means of generating a reactive intermediate, for instance, a radical ion, and as the means to monitor its reactions to products. This chapter is meant to serve both as a guide for the beginner and as an overview for the nonelectrochemist with a need to know some of the methods available. Approximately half of the chapter is concerned with various aspects of linear sweep and cyclic voltammetry in view of the importance and widespread use of these techniques. Some general aspects of the heterogeneous electron transfer process, and the chemical reactions associated with it, are introduced in this part. The reader is strongly encouraged to consult Chapter 1 in which basic electrochemical concepts are discussed in detail and monographs on the topic [1,2]. In most cases, electrons are transferred one by one [3]. Depending on the ordering of potentials of individual steps in multi-electron reactions, more complicated behavior may be observed (see, e.g., Reference 4, and the discussion in Section II.C in Chapter 11). In order to preserve the formalism used in electrochemistry, the number of electrons n, that is, transferred is maintained in most formulas although its value equals one in practically all cases. The general discussion of electrochemical reactions is, for the sake of consistency, restricted to cover reductions only. The transposition to oxidations should not present any problem, but the reader should be aware that the plus or minus sign in some equations has to be changed.

II.

LINEAR SWEEP VOLTAMMETRy AND CyCLIC VOLTAMMETRy

A. INTRODUCTION Voltammetry in an unstirred solution where the predominant mode of mass transport is limited to diffusion is one of the most useful techniques for the study of electrochemical reactions [1,2,5,6] (however, see the critical remarks in Reference 7). Most often, a triangular potential–time waveform

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Potential, E (V)

–1.2 –1.4 –1.6 –1.8 –2.0

0

5

(a)

10

15

20

Time, t (s) Ep = –1.57 V

1.2 1.0

Ep/2

0.8 i (mA)

0.6 ip

0.4 0.2

Ep = –1.22 V

ip/2

0.0 –0.2 –0.4 –1.0 (b)

–1.2

–1.4

–1.6

–1.8

–2.0

E (V)

FIgURE 2.1 (a) Potential–time waveform for a triangular voltage sweep between −1 and −2 V at a sweep rate (ν) of 0.1 V s−1 and (b) a simulated (DigiSim®) voltammogram for a substrate that is reduced at Ep = −1.57 V to a product or intermediate, the oxidation of which is observed at Ep = −1.22 V.

with equal positive and negative slopes is used with v = dE/dt being the sweep or scan rate. Usually, the initial potential (Einitial) and final potential (Efinal) are the same as illustrated in Figure 2.1a. This has given rise to the term cyclic voltammetry (CV). However, sometimes the voltage sweep is continued to include one or more additional E–t half-cycles or includes more complicated sawtooth-like waveforms to meet special needs. An example of a current–potential curve, a so-called voltammogram, resulting from a simple triangular sweep is shown in Figure 2.1b. There are two conventions to draw cyclic voltammograms: according to the American or polarographic convention, more negative potentials are plotted to the right (Figure 2.1b), while according to IUPAC (European Convention), positive potentials are shown in that direction [8]. The experiment should preferably start from an equilibrium situation, that is, close to the rest potential of the working electrode in the electrolyte, with a current i = 0. The graphical display of the currents is adjusted so that the voltammogram evolves in a clockwise direction. Thus, the two conventions are also different in the sign of the current (reductive/ cathodic or oxidative/anodic currents denoted by positive values, respectively; the former is used in Figure 2.1b). A characteristic feature is the presence of peaks, in this case two, identified by the peak potential Ep. The particular voltammogram shown in Figure 2.1b results from a process in which a substrate during the forward, negatively going voltage sweep is reduced at Ep = −1.57 V to a product or intermediate, the oxidation of which is observed at a peak potential Ep = −1.22 V during the backward (or reverse),

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positively going sweep. In addition to Ep, a voltammogram is usually characterized by the peak current ip and often also by the half-peak width Ep − Ep/2, where Ep/2 is the half-peak potential, that is, the value of E at which i = ip/2 (see Figure 2.1b). Analysis of how Ep, Ep − Ep/2, and ip vary with, for instance, changes in the voltage sweep rate, ν, and the bulk concentrations of the substrate, CA* , and a reagent, CX* , may provide information about the thermodynamics and kinetics of the followup reaction. This type of analysis is discussed in some detail in Sections II.C and II.D. If only the voltammogram corresponding to the first half-cycle is used for data analysis, the technique is usually referred to as linear sweep voltammetry (LSV). Multicycle experiments finally lead to what has been termed ultimate cyclic voltammograms [9] that represent a steady-state-like behavior with curves of subsequent cycles being superimposed. The method dates back to experiments by Randles [10a,b] and Ševčík [10c]. Significant theoretical contributions to this type of voltammetry were published in 1955 by Matsuda and Ayabe [11] and in 1964 by Nicholson and Shain [12]. The years to follow were a period of intense activity in the calculation of the electrode response for many different mechanisms. Most notable of the large volume of papers published are those of Nicholson and coworkers [13–16] and of the Savéant group [2,17–25]. Efforts initiated by Parker and coworkers [26–37] have been directed toward the development of LSV and CV as practical tools for quantitative studies of electrode processes. Nowadays, CV is probably the most popular electrochemical technique in organic electrochemistry [4–6,38–41]. One key element for this success may be the strong interaction between experimental and theoretical development, leading to user-friendly commercial simulation software (DigiSim from BAS), which is now available and has been used in this chapter for illustration purposes. The reader interested in simulation may obtain additional information about the simulation of cyclic voltammograms from Chapter 5.

B.

EXPERIMENTAL SETUP

1. The Electrochemical Cell An example of the instrumentation for LSV and CV is shown in Figure 2.2. The cell is a simple and convenient design consisting of a 30 mL tube fitted with a B19/26 ground-glass joint. The electrode holder is made of Teflon and has holes for the three electrodes [1]. In addition, there is an inlet for an inert gas, usually nitrogen or argon, by which the voltammetry solution is purged before the measurements are made. This serves to remove dissolved oxygen, which, during the study of reduction processes, may itself be reduced or may react with intermediates such as radical anions produced at the cathode. Oxygen may react also with radical cations generated at the anode. Therefore, even for oxidations, it is recommended that measurements are carried out in the absence of oxygen. The fixed arrangement of the electrodes is desirable to avoid changes in the geometry when an electrode is taken out of the electrolyte and then replaced. The working electrode usually is a circular disk made of platinum, gold, or glassy carbon, which together with the electrical connection is fitted into a nonconducting material and polished to mirror quality. For reductions, mercury (film) electrodes are frequently used also owing to their microscopic smoothness and because of the large overpotential for hydrogen evolution characteristic for this electrode material. The latter makes possible the study of difficultly reduced substrates in water and other hydroxylic solvents. The oxidation of mercury at a low potential (0.3–0.4 V vs. SCE) to mercury salts or organomercurials prevents the use of these electrodes for oxidations. Reference electrodes are frequently commercial aqueous calomel electrodes or similar. However, the use of these types of electrodes may cause unnecessary complications, such as the need to separate the reference and working compartments of the cell in order to avoid exchange of dissolved species or solvents if the electrochemical process is studied under nonaqueous conditions (see, e.g., Chapter 7). In the latter case, it is often more convenient to use nonaqueous

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Computer i

E Function generator E

Potentiostat Current–voltage converter W

R

C

t N2 Teflon stopper

Test tube

FIgURE 2.2 Experimental setup for linear sweep and cyclic voltammetry. W, working electrode; R, reference electrode; C, counter electrode.

reference electrodes most simply consisting of a silver wire immersed in the same solventsupporting electrolyte solution as that in the working compartment [42]. A disadvantage of the latter type of reference electrode is that the potential usually cannot be accurately predicted. Thus, it is necessary to calibrate the electrode, for instance, by recording the voltammogram of a standard redox couple, if the measured potentials are to be used in comparison with published values. A commonly used standard is the ferrocene/ferrocenium (Fc/Fc+) redox couple [43] (see also Chapter 7). The counter electrode usually consists of a platinum wire or a small platinum foil. Without any further refinement, the cell shown in Figure 2.2 is suitable for measurements at voltage sweep rates from about 0.1 to 500 V s−1 using working electrodes with surface diameters ranging from about 2 to 0.1 mm. Higher sweep rates will need more sophisticated instrumentation (see Section II.B.3). 2. The Solvent-Supporting Electrolyte System For routine work, solvents and supporting electrolytes of the highest commercial quality may usually be used without further purification. It is, however, often necessary to have solventsupporting electrolyte solutions that are essentially free of reactive impurities in order to study the voltammetric behavior of highly reactive intermediates. Elaborate vacuum line apparatus has been described for this purpose [5,44–47]. Sometimes, it is more convenient instead to conduct the voltammetric experiments over neutral alumina [48]. The method is simply to add active neutral alumina directly to the voltammetric cell, for example, of the type shown in Figure 2.2, and to mix vigorously for a few minutes before the measurements are made. The method is

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very effective in removing the electrophilic and nucleophilic impurities normally present in trace amounts even in the most carefully purified solvent-supporting electrolyte solutions. However, one should be aware that also the substrate may be adsorbed at the alumina, and therefore this approach should in general be used only in qualitative work where the main goal is to suppress unwanted follow-up reactions. 3. The Electronic Equipment A discussion of the instrumental aspects of voltammetry and leading references to the original literature can be found in some of the monographs already cited in Section I [1,2,5,6]. The essential units are the potentiostat, a triangular waveform generator, and a recording device, nowadays most often a digital computer. Commercial equipment with the units combined into one instrument controlled by a PC is available from a number of manufactures. Also home-built instrumentation has found use in many cases, for instance, in studies in which so-called ultramicroelectrodes (see Section III) have been used for CV at sweep rates exceeding 10,000 V s−1. Sweep rates up to 10 MV s−1 have been achieved with improved electronic circuitry aimed at minimizing artifacts in the nanosecond time scale [49]. Thus, extremely fast processes can be observed with a simultaneous extremely high spatial resolution and with diffusion layer thicknesses approaching those of the electric double layer.

C.

SIMPLE ELECTRON TRANSFER REACTIONS

In this section, we examine the relationship between current and potential in the case where the primary product of the electrode reaction is nonreactive, that is, there are no chemical reactions coupled to the electron transfer reaction at the time scale of the experiment. An electron transfer reaction, Equation 2.1, is characterized thermodynamically by the standard potential Eo, that is, the value of E at which the activities of the oxidized form A and the reduced form B of the redox couple are the same. Thus, the second term in the Nernst equation, Equation 2.2, cancels. Here, R is the molar gas constant (8.314 J K−1 mol−1), T is the temperature (K), n is the number of electrons, and F is Faraday’s constant (96,485 C). Parentheses are used for activities, brackets for concentrations, and fA and f B are the activity coefficients. However, what may be measured directly is the formal potential Eo′ defined in Equation 2.3. It follows that the relation between Eo and Eo′ is given by Equation 2.4. In this chapter, we shall assume that activity coefficients are unity and therefore that Eo′ = Eo: k

s,f o  ⇀ A + ne − ↽  B (E ) k s,b

E = Eo +

(2.1)

RT ( A) RT f [ A] ln = Eo + ln A nF (B) nF fB [ B]

(2.2)

RT [ A] ln nF [ B]

(2.3)

RT f ln A nF fB

(2.4)

E = E o′ +

E o′ = E o +

The kinetics of the electron transfer reaction, Equation 2.1, are described by the heterogeneous electron transfer rate constants ks,f and ks,b (in units of cm s−1), where the subscript s indicates a surface

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103

process. The values of ks,f and ks,b depend exponentially on the potential E as seen in Equations 2.5 and 2.6, in which ko is the standard heterogeneous electron transfer rate constant and α is the electrochemical transfer coefficient [50] (see Chapter 1). It is seen from Equations 2.5 and 2.6 that ks,f = ks,b = ko at E = Eo. The relation between ks,f and ks,b, and the current, i, is given by the Butler– Volmer equation, Equation 2.7, where A is the electrode area (cm2) and [A]s and [B]s are the surface concentrations of A and B:  −αnF ( E − E o )  ks,f = k oexp   RT  

(2.5)

 (1 − α)nF ( E − E o )  ks,b = k oexp   RT  

(2.6)

i = nFA ( ks,f [ A]s − ks,b [ B]s )   −αnF ( E − E o )   (1− α)nF ( E − E o )   [ B − = nFAk o [ A]s exp  ] exp s    RT RT      

(2.7)

In electrochemistry, electron transfer reactions are classified as reversible, quasireversible, or irreversible depending on the ability of the reaction to respond to changes in E. In voltammetry, the relevant kinetic parameters are ko, α, and ν. The mutual influence of ko and ν is conveniently expressed through the magnitude of the dimensionless parameter, Λ, defined in Equation 2.8 [11], where DA and D B are the diffusion coefficients for A and B, respectively. Usually, these are assumed to be identical, DA = D B = D. α /2

D  k  A ko  DB  Λ= = 1/ 2 1/ 2  DAνnF   DνnF   RT   RT      o

for DA = DB = D

(2.8)

1. Reversible Electron Transfer A reversible electron transfer is, strictly speaking, the limiting case where A and B are in thermodynamic equilibrium at the electrode surface, that is, the electron transfer reaction responds instantaneously to a change in E. Thus, the ratio between [A]s and [B]s depends only on E − Eo, the relationship being given by the rearranged Nernst equation:  nF ( E − E o )  [ A]s = exp   [ B]s RT  

(2.9)

The equilibrium condition implies an infinitely large value of Λ that may result from an infinitely large value of ko and/or an infinitely small value of ν. Under these conditions, the overall process is controlled by the diffusion of A and B to and from the electrode. In practical work, an electron transfer process is called reversible if the deviations from this limiting case are too small to be detected experimentally. This happens typically when Λ is larger than approximately 12 (see Section II.C.4). It also follows from the earlier discussion that the shape and the position of the voltammogram, defined, for instance, by the values of Ep, Ep − Ep/2, and ip, are independent of the parameters ko and α. A cyclic voltammogram for a reversible one-electron reduction is shown in Figure 2.3.

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Organic Electrochemistry 1.0

red

Ep/2

0.8

Epred

0.6

ipred

i (mA)

0.4

ipred/2

0.2 0.0 –0.2

ipox

–0.4 –0.6

Epox

–0.8 –1.0

–1.2

–1.4

–1.6

–1.8

–2.0

E (V)

FIgURE 2.3 Simulated (DigiSim®) cyclic voltammogram for a reversible one-electron reduction with Eo  =  −1.5 V. The other simulation parameters are ko = 104 cm s−1, α = 0.5, T = 298.2 K, C A* = 1 mM, and A  =  1  cm2. These are the same for the other simulations included in this section unless otherwise stated. The simulations do not include effects caused by uncompensated solution resistance (see Section II.E.2) and double-layer charging (see Chapter 1) or other background contributions.

Ideally, the part of the voltammogram recorded during the forward scan satisfies the following three criteria [11], where the values given in mV refer to T = 298 K: 1. The value of Ep − Eo is given by Equation 2.10 and is independent of ν. This is often expressed as in Equation 2.11: Ep − E o = −1.11

RT 28.5 V=− mV nF n

dEp =0 d log ν

(2.10)

(2.11)

2. The value of Ep − Ep/2 is given as follows: Ep − Ep/2 = −2.20

RT 56.5 V=− mV nF n

(2.12)

3. The value of ip is given by Equation 2.13, sometimes called the Randles–Ševčík equation. It is seen that ip increases linearly with ν1/2: 1/ 2

 nF  ip = 0.4463nFACA* DA1/ 2ν1/ 2    RT 

(2.13)

If also the part of the voltammogram recorded during the backward sweep is included, we have the following: 4. The peak current ratio −ipox /ipred  is unity and independent of ν. For a simple electron transfer process, measurements of the peak current ratio serve to control the assumption that B does not react on the time scale of the experiment.

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5. The peak separation, ∆Ep = Epox − Epred, is equal to approximately 57 mV, the exact value being dependent on the potential, Eswitch, at which the voltage sweep is reversed. The following equation refers to the limiting case Eo − Eswitch = ∞, where ΔEp for obvious reasons is twice the value of Ep − Eo: ∆Ep = Epox − Epred = 2 ⋅ 1.11

57.0 RT V= mV nF n

(2.14)

6. The average of Epox and Epred is often referred to as the midpoint potential E, which is approximately equal to Eo (Equation 2.15). Again, this equation is only strictly valid for Eo − Eswitch = ∞, but the error at, for instance, Eo − Eswitch = 0.3 V is negligible in most practical works:

Eo ≈ E =

(E

ox p

+Epred

)

(2.15)

2

A difficulty arises in determining −ipox /ipred in that it is not obvious how the baseline for ipox should be determined. Indeed, experimental voltammograms are often less well-behaved than the simulated one shown in Figure 2.3. The most reliable procedure appears to be a graphical method [51], which, however, should be used with care [52]. It is often possible to observe a second electron transfer reaction, Equation 2.16, within the potential window defined by the discharge of the solvent-supporting electrolyte solution. An example of a voltammogram for two consecutive one-electron reductions is shown in Figure 2.4. (2.16)

B + ne − ⇌ C

When the two electron transfer reactions are as well separated as those shown in Figure 2.4, the difference between the peak potentials for the first and second electron transfers is a good approximation to the difference in the standard potentials, E2o − E1o. When the two electron transfer reactions are closely spaced, this difference may be determined from the half-peak width of the overlapping waves [53]. However, depending on the exact value of E2o − E1o , a variety of shapes of 1.5 1.0

i (mA)

0.5 0.0 –0.5 –1.0 –1.0

–1.2

–1.4

–1.6

–1.8

–2.0

E (V)

FIgURE 2.4 Simulated (DigiSim®) cyclic voltammogram for two consecutive reversible one-electron reductions with E1o = −1.3 V and E2o = −1.7 V. For other simulation parameters, see Figure 2.3.

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the voltammograms is observed (see Chapter 11, and Reference 305.). When E2o − E1o is known, the equilibrium constant for disproportionation, Kdispr, Equation 2.17, may be calculated from Equation 2.18. Introduction of the Eo values used in the simulation for Figure 2.4, that is, E1o = −1.3 V and E2o = −1.7 V, together with n = 1 results in Kdispr = 1.7·10 –7. The value of Kdispr is dependent on solvent and is, as a rule, found to decrease with decreasing polarity of the solvent [48]. The formation of more highly charged species may be observed in special cases such as, for example, the reduction of C60 and its derivatives (see Chapter 21): 2B ⇌ A + C ( K dispr )

K dispr

(

(2.17)

)

 E2o − E1o nF   = exp  RT    

(2.18)

2. Quasireversible Electron Transfer In the general case, named quasireversible, the electron transfer reaction, Equation 2.1, does not respond instantaneously to changes in E. In other words, [A]s and [B]s are determined not only by the value of E − Eo, but also, via Equations 2.5 and 2.6, by the magnitudes of ko and α. There is a mixed control by diffusion and electron transfer kinetics. Typical voltammograms for quasireversible electron transfers are shown in Figure 2.5. In comparison with the voltammogram for the reversible case (Figure 2.3), it is seen that the reduction peak has moved in the negative direction and the oxidation peak in the positive direction resulting in a peak separation ΔEp larger than ~57 mV. It appears from Equations 2.5 and 2.6 that decreasing values of α cause ks,f to decrease and ks,b to increase, and as a consequence, Epred and Epox both move in the negative direction (Figure 2.5). Conversely, increasing values of α cause Epred and Epox to move in the positive direction (Figure 2.5). The overall result is that the effect of the magnitude of α on ΔEp is small as long as α does not deviate too much from 0.5 [13,54]. In that case, ko may conveniently be determined from values of ΔEp recorded at different values of ν. It is important to notice, however, that the mere observation of ΔEp larger than ~57 mV (at T = 298 K) is not in itself a sufficient criterion for the classification of an electron transfer reaction as quasireversible. Many 0.8 0.6 0.4 i (mA)

0.2 0.0 –0.2 –0.4 –0.6 –1.0

–1.2

–1.4

–1.6

–1.8

–2.0

E (V)

FIgURE 2.5 Simulated (DigiSim®) cyclic voltammograms at ν = 1 V s−1 for quasireversible one-electron reductions with Eo = −1.5 V, ko = 3·10 −3 cm s−1, and α = 0.3 (dot), α = 0.5 (full), and α = 0.7 (dash). For other simulation parameters, see Figure 2.3.

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other factors, such as adsorption phenomena and insufficient electronic compensation of the voltage drop caused by the solution resistance (see Section II.E), may cause voltammograms to have peak separations larger than ~57 mV. The expressions for Ep − Eo, Ep − Ep/2, and ip for the quasireversible case are given in Equations 2.19 through 2.21, where Ξ(Λ, α), Δ(Λ, α), and Κ(Λ, α) are nonlinear functions of Λ and α. These are available as graphical representations in the paper by Matsuda and Ayabe [11]. Alternatively, the reader may find expressions for Ep − Eo, Ep − Ep/2, and ip in dimensionless notation in a paper by Nadjo and Savéant [19]. Ep − E o = −Ξ(Λ, α)

RT nF

(2.19)

Ep − Ep/2 = −∆(Λ, α)

RT nF

(2.20)

1/ 2

 nF  ip = 0.4463Κ (Λ, α)nFACA* DA1/ 2ν1/ 2    RT 

(2.21)

3. Irreversible Electron Transfer When Λ becomes progressively smaller, the shape of the voltammogram continues to change. Experimentally, a constant shape is reached when Λ is smaller than approximately 0.2 (see Section II.C.4). The value of E required to obtain an appreciable rate of reduction of A is now so much on the negative side of Eo that the second term in Equation 2.7 may be neglected. In other words, the electron transfer has become irreversible. By the same type of argument, it is clear that the oxidation of B back to A during the backward sweep proceeds irreversibly as well. Sometimes this is called electrochemically irreversible to make a distinction to chemically irreversible (irreversible follow-up reaction). Typical voltammograms for irreversible electron transfers are shown in Figure 2.6. In comparison with the voltammograms for the quasireversible case (Figure 2.5), it is seen that the reduction peak has moved even more in the negative direction and the oxidation peak even more in the positive direction, now to the point that there is a potential region between the two peaks in

0.8 0.6

i (mA)

0.4 0.2 0.0 –0.2 –0.4 –0.5

–1.0

–1.5

–2.0

–2.5

E (V)

FIgURE 2.6 Simulated (DigiSim®) cyclic voltammogram at ν = 1 V s−1 for irreversible one-electron reductions with Eo = −1.5 V, ko = 10 −5 cm s−1, and α = 0.3 (dot), α = 0.5 (full), and α = 0.7 (dash). For other simulation parameters, see Figure 2.3.

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which essentially no current flows. The extension of this region depends on the magnitude of Λ. Also, the effect of α is more pronounced than for the quasireversible case. Ideally, the voltammogram for the irreversible case satisfies the following three criteria [11], where, as before, the values given in mV refer to T = 298 K: 1. The value of Ep − Eo is given by Equation 2.22 and depends on α, ko, DA, and ν. It follows that Ep changes with ν as given in Equation 2.23:  ko 1 ανnF  RT Ep − E o =  −0.783 + ln 1/ 2 − ln  DA RT  αnF 2 

(2.22)

dEp 1 RT 29.6 =− ln 10 = − mV 2 αnF αn d log ν

(2.23)

2. The value of Ep − Ep/2 is given in the following equation, which for α = 0.5 results in Ep − Ep/2 equal to −95.4 mV: Ep − Ep/2 = −1.857

RT 47.7 V=− mV αnF αn

(2.24)

3. The value of ip is given by Equation 2.25. It is seen that ip increases linearly with ν1/2. By comparing Equations 2.13 and 2.25, it is seen that ip for the irreversible case is 1.11α1/2 (= 0.78 for α = 0.5) times ip for the reversible case. 1/ 2

 αnF  ip = 0.4958nFACA* DA1/ 2ν1/ 2    RT 

(2.25)

4. Concluding Remarks and Examples The borderlines between reversible, quasireversible, and irreversible behavior were originally defined by Matsuda and Ayabe [11] on the basis of mathematical reasoning. However, in practical work, it is more convenient to define borderlines reflecting where deviations from the two limiting cases, reversible and irreversible, may be observed experimentally. Using the conservative estimate (see Section II.E.3) that the error in peak potential measurements is typically ±2 mV, Nadjo and Savéant [19] arrived at the borderlines given by Equations 2.26 through 2.28 for α = 0.5. The handy expressions for ko refer to n = 1, T = 298 K, and D = 1·10 −5 cm2 s−1, and assume that ν is measured in V s−1: Reversible: Λ ≥ 11.5 or k o ≥ 0.23ν1/ 2 cm s−1

(2.26)

Quasi-reversible:11.5 ≥ Λ ≥ 0.2 or 0.23ν1/ 2 cm s−1 ≥ k o ≥ 0.004ν1/ 2 cm s−1

(2.27)

Irreversible: Λ < 0.2 or k o < 0.004ν1/ 2 cm s−1

(2.28)

It follows from the earlier discussion that an electron transfer reaction that appears reversible at one (low) sweep rate may change to a quasireversible or even an irreversible process at higher sweep rates. This should be kept in mind since the application of LSV and CV in kinetics and mechanism studies, detailed in Section II.D, includes the recording of voltammograms at a number of different sweep rates, and the data analysis is usually based on the assumption that the electron transfer is reversible.

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109

Radical ions and radicals are usually highly reactive and react further on the time scale of voltammetric experiments at low ν. Exceptions are many heterocyclic compounds and substrates carrying a substituent that is able to stabilize a negative or positive charge and/or an unpaired electron. Examples include the reduction of nitro compounds (Chapter 30), quinones (Chapter 26), and compounds such as viologens containing two reducible functional groups [55] or the oxidation of thianthrene (Chapter 27), tetrathiafulvalenes (Chapter 27), and a number of methoxy-substituted aromatic hydrocarbons (Chapter 18). Numerous other examples may be found throughout this book. In many of the examples mentioned, it is possible to observe also the second electron transfer leading to the dianion or dication. In a special case, reversible or quasireversible electron transfers leading ultimately to the formation of an octaanion have been observed [56]. The determination of Eo for the oxidation or reduction of a substrate provides a direct measure of the free energy of formation of the resulting intermediate, a radical ion or radical. Values of Eo are important quantities in thermochemical calculations of, for instance, bond energies [57]. The temperature dependence of Eo that gives insight into ΔS for the electron transfer reaction has been investigated in a number of cases, in particular for the reduction of aromatic nitro compounds [58]. Values of −ΔS273.2 are typically in the range 5–20 cal K−1 mol−1. Pressure-dependent CV revealed the activation volume of electrode processes [59] and led to the conclusion that the electron transfer for the decamethyl ferrocene/ferricinum redox couple is controlled by solvent dynamics. Electron transfer reactions that are accompanied by large structural changes may give rise to the unusual observation that the second electron is easier to add or remove than the first, so-called potential inversion [4,60–62] (see also Chapter 11). Examples of the application of CV for the study of the kinetics of the heterogeneous electron transfer reaction include the reduction of quinones [63,64] and nitroalkanes [64–66]. For radical anions, the effect of the nature of the counterion, that is, the cation of the supporting electrolyte, has been investigated [64,67], and the general trend is that the value of ko decreases when the length of the carbon chain in R4N+ increases (see also Section II.D.5 for other counterion effects). Attention should be paid also to the fact that ko is sometimes observed to depend on the electrode material and thus is not a true standard value [68]. Subtle details of electron transfer reactions can be revealed, such as the difference between inner- and outer-sphere electrode reactions [69] or the type of kinetic relationship (Butler–Volmer versus Marcus–Hush) [70]. In addition to CV, methods based on AC voltammetry are useful for studies of heterogeneous electron transfer kinetics [71]. This technique where a sinusoidal waveform is superimposed on the classical cyclic voltammetric triangle can be viewed as a special case of voltammetries with any periodic potential applied to an electrode [72a,b]. The experimental current variation with time is advantageously analyzed by applying a Fourier transformation (FT-CV), resulting in a power spectrum in the frequency domain. Selection of particular frequency bands and inverse Fourier transformation generates not only the classical (DC) voltammogram but also fundamental, second-, third-, and higher harmonic signals. The full potential of the technique is revealed by comparison to simulated curves [72c,d] and reveals particular patterns depending on the prevalent electrode reaction mechanisms [72e]. The analysis was extended to coupled chemical reactions [72f,g] and electron transfer of surfaceconfined proteins [72h,i]. As an additional feature, separation of faradaic and background contributions to the currents was achieved [72j]. The FT technique was also applied to experiments at the rotating disk electrode [72k] (RDE; see Section VI.B), allowing separation of transport modes in the harmonic analysis, thus providing a full characterization of electron transfers. Although FT-CV might not be readily available in most electrochemical research labs, it might become a valuable tool in the future owing to the wealth of information provided. Recent examples are the observation of dianions derived from trans-stilbenes [73], the study of tetrathiafulvalene oxidation electrode kinetics [74], and the analysis of the redox behavior of flavin adenine dinucleotide in the redox center of glucose oxidase [75].

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D.

Organic Electrochemistry

ELECTRON TRANSFER REACTIONS FOLLOWED BY CHEMICAL REACTIONS IN SOLUTION

Some of the most successful applications of LSV and CV are concerned with the study of the kinetics and mechanisms of the reactions of electrode-generated intermediates, and a large share of the classical electrochemical literature in this field deals with this aspect of voltammetry [12–37,76–78]. The majority of electrochemical reactions include radical ions as the primary intermediates, and the reaction schemes describing the conversion of a substrate A to products are typically composed of one or two one-electron transfers and one or two chemical steps. Examples include hydrogenations, (+2e−, +2H+) (see Chapters 22 and 44) and hydrodimerizations (+e−, +H+) (see Chapter 17), and anodic additions (−2e−, +2Nu−) (see Chapter 19), dehydrodimerizations (−e−, −H+) (see Chapter 18), and substitutions (−2e−, +Nu−, −H+), where Nu− is a nucleophile (see Chapter 19). The electrochemical literature abounds with symbols and abbreviations that are not always strictly logical. Therefore, a few comments about the abbreviations used in the following to designate basic electrode mechanisms are necessary. The notation is based on that due to Testa and Reinmuth [79], where the letter e indicates an electron transfer process and the letter c indicates a chemical reaction. It is helpful to distinguish between reversible (fast) and rate-determining (slow) steps by using lowercase letters for reversible and capital letters for rate-determining steps. Since the second electron transfer reaction can take place either at an electrode (heterogeneous electron transfer) or in solution (homogeneous electron transfer, that is, in the diffusion layer or in the bulk solution), those taking place in solution are given the subscript h. The use of this and other abbreviations frequently met in the literature is illustrated in the following. The presentation is restricted to cover only chemical reactions that follow a reversible electron transfer. Reaction schemes including quasireversible electron transfer or dissociative electron transfer reactions [80,81], for instance, observed for the reduction of alkyl halides (Scheme 2.1), are mentioned only briefly (see Chapters 14, 24, and 25 for details). 1. Typical Irreversible Follow-Up Reactions Examples of mechanisms resulting in first-order rate laws are summarized in Scheme 2.2, where the right-hand part of (ii) refers to the situation where the reaction of B is with a reagent X. In that case, R-X + e–

SCHEME 2.1

Dissociative electron transfer reaction.

A + e– kii B

(i)

B C

or

B+X

kii'

C

(ii)

C + e–

D

(iii)

B+C

A+D

(iv)

D+X

E

(v)

Mechanism abbreviation

Observed rate law

kobs (CX* /CA* >>1)

(i)–(ii):

eC

–d[B]/dt = kobs[B]

kii or kii'CX*

(vi)

(i)–(ii)–(iii): eCe

–d[B]/dt = kobs[B]

kii or kii'CX*

(vii)

(i)–(ii)–(iv): eCeh

SCHEME 2.2

R + X–

–d[B]/dt = kobs[B]

Mechanisms resulting in first-order rate laws.

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2kii or

2kii'CX*

(viii)

111

Techniques for Studies of Electrochemical Reactions in Solution A + e– 2B A+B

SCHEME 2.3

(i)

B kii kiii

C

(ii) (iii)

I

B+I

A+C

(iv)

C+X

D

(v)

Mechanism abbreviation

Observed rate law

kobs

(i)–(ii): eC(dim) or RR

–d[B]/dt = kobs[B]2

2kii

(vi)

(i)–(iii)–(iv): RS

–d[B]/dt = kobs[A][B]

2kiii

(vii)

Mechanisms resulting in second-order rate laws.

the kinetics are often studied under pseudo-first-order conditions, that is, at CX∗ /CA∗ ≫ 1. Often, the product C is more easily reduced than A [82] resulting in a second electron transfer. This may take place either at the electrode (iii) or in solution (iv) [83] resulting in an eCe or an eCeh mechanism, the latter sometimes being referred to as a DISP1 mechanism [84,85]. The intermediate D finally reacts with, for example, the reagent X, to the product E. The observed rate laws (vi) through (viii) for all these cases are the same if step (ii) is rate determining. Another important set of reactions are the dimerizations, which are usually discussed within the frames of the two mechanisms shown in Scheme 2.3. If the formation of C results from the coupling of two radicals or radical ions (ii), the reaction is referred to as a radical–radical (RR) process. Another route to C, the radical–substrate (RS) process, includes the coupling between A and B (iii) resulting in the formation of an intermediate I that is further reduced to C by reaction with B (iv). The direct reduction of I to C at the electrode is usually without importance. Finally, the intermediate C reacts with a reagent X to the product D. These dimerization mechanisms belong to a more general scheme to be discussed in some detail later (Section II.D.5). 2. Kinetic Classification of Simple Irreversible Follow-Up Reactions Let us now consider the voltammetric response for a reversible one-electron reduction followed by an irreversible chemical reaction, for instance, the eC mechanism with the rate constant k (= kii in Scheme 2.2). The voltammograms resulting from different values of k at ν = 1 V s−1 are shown in Figure 2.7. Given the value of ν, it is seen that both the shape and the position of the voltammogram depend on the magnitude of k. On the other hand, given the value of k, it is intuitively understood that the effect of the chemical reaction will gradually diminish if the sweep rate is allowed to increase. In the limit, the experiment time is so short that the chemical reaction does not have the time to manifest itself, and consequently, the voltammogram observed is just that for the one-electron transfer reaction, Equation 2.1. The effect of the relative magnitudes of k and ν on the position and shape of a voltammogram is conveniently discussed in terms of the dimensionless parameter λ, defined by Equation 2.29 for a reaction following a first-order rate law and by Equation 2.30 for a reaction following a second-order rate law [17–19]. λ=

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kRT νnF

(firstorder rate law)

(2.29)

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Organic Electrochemistry 1.0 0.8 0.6 i (mA)

0.4 0.2 0.0 –0.2 –0.4 –0.6 –0.8 –1.0

–1.2

–1.6

–1.4

–1.8

–2.0

E (V)

FIgURE 2.7 Simulated (DigiSim®) cyclic voltammogram at ν = 1 V s−1 for an eC mechanism with Eo = −1.5 V and k = 2 s−1 (dash), 10 s−1 (dot), 102 s−1 (dash-dot), 10 4 s−1 (dash-dot-dot), 10 6 s−1 (short dash), and 108 s−1 (short dot) corresponding to λ = 0.0513, 0.257, 2.57, 2.57·102, 2.57·104, and 2.57·106, respectively. The full line corresponds to the simple electron transfer reaction shown in Figure 2.3, where also other simulation parameters are given.

λ=

kCA* RT νnF

(secondorder rate law)

(2.30)

It is convenient to classify the electrochemical reactions with regard to the nature of the response. Given the value of ν, it is seen from Equations 2.29 and 2.30 that increasing values of k correspond to increasing values of λ. Beginning with λ = 0, the effect of increasing values of λ is initially that the current associated with the oxidation of B during the backward sweep gradually disappears (see, e.g., Figure 2.7). The peak potential is, so far, only little affected. The region of λ values, from the point where the effect of the chemical reaction becomes experimentally detectable, to the point where the current for the oxidation of B has totally disappeared is given by Equation 2.31 for the eC mechanism (first-order rate law) and Equation 2.32 for the RR dimerization (second-order rate law) [19]. Again, the handy expressions correspond to n = 1 and T = 298 K. In this region, the system is under mixed diffusion and kinetic control. Limits for λ relating to other mechanisms, including the eCe, eCeh, and more complex reaction schemes, have been reported in the literature [17–19]: eC mechanism: 0.11 < λ < 1.89 or 4.3ν s−1 < k < 73.5ν s−1

(2.31)

RR dimerization: 0.37 < λ < 1.35 or 14.5ν s−1 < kCA* < 52.5ν s−1

(2.32)

At λ values higher than those given, for instance, by the upper limits in Equations 2.31 and 2.32, the shape of the voltammogram is essentially independent of λ, and an increase in λ only results in a displacement of the voltammogram in the positive direction. This is illustrated by the voltammograms corresponding to k = 104, 106, and 108 s−1 in Figure 2.7. A stationary state has now been established in solution by mutual compensation of the chemical reaction of B and the diffusion process, and the system is said to be under purely kinetic control. The limits given by Equation 2.31 or 2.32 define zones of particularly controlled regimes. Note that for more complex mechanisms, such an approach yields 2D or 3D zone diagrams [2] (see also Chapters 5 and 10). The peak current in the first sweep is only slightly affected by the value of λ, with an increase at small λ to a somewhat larger value as compared with the reversible case.

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3. Irreversible Follow-Up Reactions, Mixed Diffusion, and Kinetic Control (CV and DCV) As mentioned earlier, the characteristic features of processes in this category are that Ep is close to that for the no-reaction case, Equation 2.1, and that the peak current ratio −ipox /ipred varies from approximately unity to zero. The observation of oxidation current for B during the backward sweep shows that the material conversion is low. By comparison of the voltammograms for the eC and the eCeh mechanisms in Figure 2.8, it is seen that the second electron transfer reaction in the eCeh mechanism gives rise to only little additional current illustrating that only a small fraction of B has been converted to C. Although the value of E p in this kinetic region is only slightly affected by the magnitude of λ, a  careful investigation of the sweep rate dependence of Ep may, nevertheless, have diagnostic value, for instance, in the distinction between the RR- and RS-dimerization mechanisms. This is illustrated in Figure 2.9, which shows the theoretical curves, the so-called working curves, for 1.2 1.0 0.8 i (mA)

0.6 0.4 0.2 0.0 –0.2 –0.4 –1.0

–1.2

–1.4

–1.6 E (V)

–1.8

–2.0

FIgURE 2.8 Simulated (DigiSim®) cyclic voltammograms for the eC mechanism (full line) and the eCeh mechanism (dot) at ν = 1 V s−1 with Eo = −1.5 V and k = 4 s−1 (corresponding to λ = 0.103). For other simulation parameters, see Figure 2.3. 0

Ep– E o (mV)

a

20

40

–3

–1

1 log(λ)

FIgURE 2.9 Working curves for RR- (full line) and RS-dimerization (dotted line) mechanisms showing the predicted variation of Ep − Eo with log λ together with data obtained by LSV for the dimerization of the 7,12-diphenylbenzo[k]fluoranthene radical cation in benzene/MeCN (1/1); CA* = 0.44 mM (◊),CA* = 1.04 mM (Δ), and CA* = 2.00 mM ( ). (Reprinted with permission from Debad, J.D., Morris, J.C., Magnus, P., and Bard, A.J., J. Org. Chem., 62, 530–537. Copyright (1997) American Chemical Society.)

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Ep − Eo for the two mechanisms together with experimental data obtained for the dimerization of 7,12-diphenylbenzo[k]fluoranthene radical cation [86]. The data clearly indicate that the dimerization is of the RR type. A common procedure for studying the kinetics of follow-up reactions by CV is to compare values of −ipox /ipred recorded at different sweep rates to the working curve for the proposed mechanism [12]. However, a problem with this approach is the difficulty already mentioned in evaluating the baseline for the backward sweep, and therefore, CV suffers from some limitations when used in quantitative work. Simulation (see Chapter 5) is a common helpful approach nowadays, but originally, derivative cyclic voltammetry (DCV) was developed to solve this problem [37,77,87]. The latter has been applied to settle many mechanistic questions and has various advantages (see later in this section). The application of the first derivative of the CV curve was investigated during the 1960s by Perone and coworkers [88,89], but due to difficulties in performing differentiation using analog electronic equipment, the method was not recommended for the measurement of electrode potentials. Digital differentiation may now be carried out by using Fast Fourier Transform methods [90], and differentiation algorithms are included in modern software. The voltammogram shown in Figure 2.3 is redrawn in Figure 2.10a, now showing the current i as a function of the time t. The problem of defining a baseline for the measurement of ipox is illustrated by the two broken lines, a problem that becomes more serious as Eswitch comes closer to Eo. The differentiated curve, di/dt versus t, is shown in Figure 2.10b. The peaks labeled if′ and ib′ , corresponding to the maximum steepness of the voltammogram during the forward and backward 1.0 0.8 0.6 0.4 i (mA)

0.2 0.0 –0.2 –0.4 –0.6 –0.8 0.0

0.5

1.0

(a)

1.5

2.0

t (s) 0.06 0.04 i΄f

di/dt

0.02 0.00 –0.02

i΄b –0.04 –0.06 0.0 (b)

0.5

1.0

1.5

2.0

t (s)

FIgURE 2.10 The cyclic voltammogram in Figure 2.3 shown as (a) the current–time curve and (b) the differentiated current–time curve.

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sweeps, respectively, reflect the CV peak heights. Both if′ and ib′ are easily being measured relative to the zero line. Since the slope of the baseline for ipox, that is the extension of the reduction wave, is not far from zero where the measurement is made, it follows that the magnitude of ib′ is essentially baseline independent. Another advantage of recording the derivative signal is that the effect of the background current is strongly diminished. The ratio RI′ = −ib′ /if′ plays the same role in DCV as the ratio −ipox /ipred in CV. It should be noticed, however, that −ib′ /if′, in contrast to −ipox /ipred, does not approach zero for increasing values of λ, but a value close to 0.1. This is related to the fact that the derivative curve during the backward sweep is not zero, even when the peak owing to B has completely vanished. The kinetic analysis of a follow-up reaction by DCV involves the recording of RI′ at different ν, and the rate constant k (or kobs) is then obtained by fitting the working curve for the appropriate rate law, usually in the form of RI′ versus log λ, to these data. An example is shown in Figure 2.11, where the working curve for the RR-dimerization mechanism is fitted to the experimental data for the dimerization of (−)-bornyl cinnamate radical anions [91]. The rate constant is obtained by matching the two scales; in the present case, this results in k = 5.6·102 M−1 s−1. The approach works equally well for reactions taking place after the second electron transfer, for instance, at the dication [92] or dianion [93] oxidation state. The rate law necessary for making a mechanism suggestion is conveniently determined by DCV. The procedure, sometimes referred to as the reaction order approach, is based on measurements of the sweep rate necessary to maintain a certain constant conversion of B as a function of CA∗ , and CX∗ if relevant. Usually, the sweep rate necessary to keep RI′ equal to 0.5 is used. This sweep rate is referred to as ν1/2 or ν0.5. The reciprocal value, 1/ν1/2, is conceptually related to the half-life time t1/2 in conventional kinetics. If the generalized rate law for the reaction to be investigated is written as Equation 2.33, where kobs is defined, for instance, as in Schemes 2.2 and 2.3, the relations between ν1/2 and the reaction orders a, b, and x are given by Equations 2.34 and 2.35, where RA/B is equal to a + b and RX is equal to x. It is seen that RA (= a) and RB (= b) are not directly separable, a common feature of all reversal techniques, where A and B are in thermodynamic equilibrium at the electrode surface. a

b

Rate = kobs  A   B   X 

x

(2.33)

log ν (V s–1) 1

–1

0

1

2

0.8

RI

0.6 0.4 0.2 0 –1

0

1 –log(kCRT/nFv)

2

3

FIgURE 2.11 DCV working curve for the RR-dimerization mechanism (bottom scale) together with experimental data obtained for the dimerization of (−)-bornyl cinnamate radical anion in DMF (top scale). (From Amatore, C. and Savéant, J.M., J. Electroanal. Chem., 85, 27, 1977.)

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log (ν1/2/Vs–1)

2.4

1.6

0.8

Slope: 1.06

0.0 0.6

1.0

1.4

1.8

2.2

2.6

log (C oHB/mM)

FIgURE 2.12 Data for log ν1/2 as a function of log(CXo /mM) obtained by DCV for the protonation of the anthracene radical anion by phenol (HB) in DMSO. (From Nielsen, M.F. and Hammerich, O., Acta Chem. Scand., 43, 269, 1989. With permission.)

RA/B = 1 +

RX =

d log ν1/ 2 d log CA*

(2.34)

d log ν1/ 2 d log CX*

(2.35)

Usually, the data treatment includes that values of log ν1/2 are plotted against log CA* or log CX* , which, for a simple rate law such as Equation 2.33, results in a straight line with the slope RA/B − 1 or RX. Data obtained for the protonation of the anthracene radical anion by phenol in DMSO [94] are shown in Figure 2.12 as an example. The slope of the regression line in this case is 1.06 indicating that the rate law is first order in the concentration of phenol. Once the rate law is known, the value of k (or kobs) may in principle be obtained directly from ν1/2 as illustrated by the relations given in Equation 2.36 for the eC mechanism (first-order rate law) and Equation 2.37 for the RR dimerization (second-order rate law), in both cases for Eo − Eswitch = 0.3 V [95]. Relations for other mechanisms may be found in the literature [95]. eC mechanism: k = 0.078

ν1/ 2 nF RT

RR dimerization: k = 0.117

ν1/ 2 nF CA* RT

(2.36)

(2.37)

However, in general, it is recommended that the full working curve is used for the determination of the rate constant in view of the larger number of data points included. This approach also offers the additional benefit that the validity of the mechanism hypothesis is tested by the goodness of the fit for each determination. This is particularly important in cases such as those where changes in the substrate substitution pattern or in the experimental conditions may result in a change in the mechanism or, for complex reaction schemes, a change in the rate-determining step (see also Section II.D.5).

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The kinetic window for CV and DCV may be derived from the limits for λ given, for example, in Equations 2.31 and 2.32. Assuming that the applicable range of useful ν values is 0.1–500 V s−1, the limits translate approximately as follows: eC mechanism: 0.4 s−1 < k < 4 ⋅ 10 4 s−1

(2.38)

RR dimerization: 1.5 s−1 < kCA* < 3 ⋅ 10 4 s−1

(2.39)

Typical examples of the application of DCV for kinetic and mechanistic studies include the cleavage of the carbon–halogen bond in the radical anions of aromatic halides [96], the protonation of radical anions [94,97–99], and the dimerization of radical ions [91]. 4. Irreversible Follow-Up Reactions, Purely Kinetic Control (LSV) The characteristic feature of the voltammogram in this region of λ values is the absence of a peak during the backward sweep, which for the same reason is not included in the data analysis (LSV). The position of the voltammogram, as measured by Ep, is displaced in the positive direction relative to that for the no-reaction case, Equation 2.1, and depends markedly on the value of λ. In contrast, the shape of the voltammogram, as measured by Ep − Ep/2, is nearly constant and independent of λ. However, Ep − Ep/2 depends on the mechanism, which is of diagnostic value (see later in this section). It is appropriate to mention at this place that processes belonging to this category are often referred to as irreversible. This may be confusing, since the term irreversible in this case refers to the effect of an irreversible chemical reaction on the total process and not just to the electron transfer reaction. Common to the two types of irreversible processes is, however, the absence of an oxidation peak owing to B in the potential region where the reduction of A takes place. Sometimes the term chemically irreversible is used for the case discussed here. A slightly different meaning is adopted for this in Section III.E.2 in Chapter 1. The lack of an oxidation peak for B indicates a high material conversion, which can also be seen by comparison of the voltammograms for the eC and eCeh mechanisms in Figure 2.13. In contrast to the voltammograms in Figure 2.8, it is seen that ip for the eCeh mechanism is now close to being twice that for the eC mechanism illustrating that the conversion of B to C has proceeded almost to completion. 2.0 1.6

i (mA)

1.2 0.8 0.4 0.0 –1.0

–1.2

–1.4

–1.6

–1.8

–2.0

E (V)

FIgURE 2.13 Simulated (DigiSim®) cyclic voltammograms for the eC mechanism (full line) and the eCeh mechanism (dot) at ν = 1 V s−1 with Eo = −1.5 V and k = 106 s−1 (corresponding to λ = 2.57·104). For other simulation parameters, see Figure 2.3.

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LSV is a powerful tool for the study of processes under purely kinetic control. Theoretical analyses of the response for various mechanisms have been carried out [12,14–25], and a series of papers [35,78,87,100] has been devoted to assimilating the theoretical results in a form useful to the experimentalist. For the general rate law, Equation 2.33, the dependence of Ep on changes in log ν, log CA* , and log CX* , respectively, is linear with the slopes given by Equations 2.40 through 2.42, where a, b, and x are the reaction orders: dEp 1 RT =− ln 10 d log ν b + 1 nF

(2.40)

dEp a + b − 1 RT = ln 10 d log CA* b + 1 nF

(2.41)

dEp x RT = ln 10 d log CX* b + 1 nF

(2.42)

Introduction of the reaction orders for, for instance, the eC mechanism under pseudo-first-order conditions (a = 0, b = 1, and x = 1) results in dEp/d logν = −29.6 mV, dEp / log CA* = 0 mV , and dEp / log CX* = 29.6 mV at n = 1 and T = 298 K. Important is also the shape of the voltammogram represented by the value of Ep − Ep/2. The limiting values for the eC and RR-dimerization mechanisms are given in Equations 2.43 and 2.44, where, again, the values given in mV refer to T = 298 K. Values of Ep − Ep/2 for other mechanisms are available in the literature [17–19]. eC mechanism: Ep − Ep/2 = −1.86

RT 47.8 V=− mV nF n

RR dimerization: Ep − Ep/2 = −1.51

RT 38.8 V=− mV nF n

(2.43)

(2.44)

When Eo for the initial electron transfer reaction is known, the measurements of Ep directly gives access to the rate constant k. Examples of the relationship between Ep − Eo and k are given by Equations 2.45 and 2.46 [19], where Equation 2.46 refers to rate law (vi) in Scheme 2.3, that is, the stoichiometric coefficient 2 is not included in the rate constant. 1 kRT   RT   eC mechanism: Ep − E o =  −0.783 + ln ⋅ 2 νnF   nF    1 4kCA* RT RR dimerization: Ep − E o =  −0.902 + ln 3 3νnF 

  RT  ⋅    nF 

(2.45)

(2.46)

It is easily seen that the application of equations such as Equations 2.45 and 2.46 requires data for Ep − Eo of high precision. For example, for the RR-dimerization mechanism, a change in k by a factor of 10 corresponds only to a change in Ep of 19.7 mV (at T = 298 K). Taking into account that both the precision and the accuracy of literature values for Eo may not be better than ±10 mV, it is strongly recommended that the value of Eo to be used is the result of a measurement made in conjunction with the Ep measurements. Since, by definition, the value of λ is rather large for Equations 2.45 and 2.46 to apply, it follows that large sweep rates are required to outrun the chemical reaction.

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This is conveniently achieved by using ultramicroelectrodes for the measurements [101,102] (see Section III). Reliable Eo data may also be predicted within a series of compounds from correlations to data obtained by NMR spectroscopy [103]. Although LSV is experimentally easy to use, the technique should be used with care. The major concern is to make sure that the electron transfer reaction is reversible and that the process under investigation is indeed under purely kinetic control in the range of ν and concentrations employed. If increasing sweep rates bring the electron transfer into the quasireversible region (see Section II.D.8), the values of, for instance, dEp/d log ν, will be larger than predicted for the reversible case [19,104], and if the follow-up reaction is brought into the region of mixed kinetic and diffusion control, dEp/d log ν will be smaller than predicted [19,105]. This may lead to a wrong mechanism assignment. However, in both cases, the Ep − Ep/2 will be larger than predicted and therefore serves as an excellent diagnostic tool for probing whether the reaction being studied does indeed confine to the purely kinetic conditions as defined earlier. Another problem may arise in the study of processes of the type where B reacts with an added reagent X. In that case, the kinetic measurements are often carried out under pseudo-first-order conditions, that is, at CX* /CA* ≫ 1, and values of dEp / log CX* are obtained by carrying out a series of E p measurements at increasing values of CX* . If now the ratio CX* /CA* at the lowest concentration of CX* is not sufficiently large, the consumption of X cannot be neglected, and accordingly, the rate of the chemical reaction is too low. The problem gradually disappears with increasing values of CX* , and the overall effect is that the increase in the reaction rate with increasing CX* becomes too large, and accordingly, the resulting value of dEp / log CX* becomes too large as well [105]. This may lead to the erroneous interpretation that the reaction order in X is larger than one. Illustrations of the application of LSV for studies of reactions under purely kinetic conditions include the oxidation of 9-substituted fluorenide ions [106] and the reduction of 2,6-diphenylpyrylium ions [49a], in both cases leading to the neutral radical that dimerizes in an RR-type reaction, the oxidation of 1,4-dithiafulvenes into tetrathiafulvalenes [107], the oxidative ring opening of arylcyclopropanes [108], the reduction of fluoroalkoxyarenes in liquid ammonia [109], and the competition between protonation and dimerization during the reduction of cinnamic acid esters in MeOH [110]. 5. Reversible Follow-Up Reactions Let us now consider a reversible electron transfer reaction followed by a reversible chemical reaction. Typical examples, including the reaction of B with a reagent X (ii), here with p equivalents, and the reversible dimerization (iii), are shown in Scheme 2.4. If the chemical reaction responds instantaneously to changes in the concentration of B or C, it is adequately described by the magnitude of the equilibrium constant K, which means that the total system is reversible. For reaction (ii), the resulting voltammogram, shown in Figure 2.14a for p = 1 and CX* /CA* = 100, has the same shape as that in Figure 2.3, but is displaced on the voltage axis according to Equation 2.47, where K corresponds to K ii in Scheme 2.4. It is seen that E p changes by 59.1 mV (at T = 298 K) in the positive direction for a 10-fold increase in K for K ≫ 1. (See, e.g., the voltammograms shown in Figure 2.14a for K = 100 and 1000.) It follows from A + e–

B

B + pX

C

2B

SCHEME 2.4

C

(E 0) Kii = Kiii =

(i) [C] [B][X]p [C] [B]2

Mechanisms including reversible follow-up reactions.

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(ii)

(iii)

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Organic Electrochemistry 1.0 0.8 0.6 0.4 i (mA)

0.2 0.0 –0.2 –0.4 –0.6 –0.8 –1.0

–1.2

–1.4

(a)

–1.6

–1.8

–2.0

E (V) 3e–5 e

2e–5

a e d c

b

a

1e–5 0e+0 a = DMOBQ b = 0.06 M EtOH c = 0.2 M EtOH d = 0.47 M EtOH e = 0.9 M EtOH

–1e–5 –2e–5 –3e–5 (b)

0.0 –0.2 –0.4 –0.6 –0.8 –1.0 –1.2 –1.4 –1.6 –1.8 –2.0 Potential (V vs. SCE)

FIgURE 2.14 (a) Simulated (DigiSim®) cyclic voltammograms for a reversible one-electron reduction followed by a reversible chemical reaction (Scheme 2.4) with Eo = −1.5 V, C A* = 1 mM, CX* = 100 mM, and K = 10 (dash), 100 (dot), and 1000 (dash-dot). The full line corresponds to the simple electron transfer reaction shown in Figure 2.3, where also other simulation parameters are given. (b) Experimental cyclic voltammograms for the reduction of 2,5-dimethoxy-1,4-benzoquinone (DMOBQ) in benzonitrile containing Bu4NPF6 (0.1 M) and different concentrations of ethanol. (Reprinted with permission from Gupta, N. and Linschitz, H., J. Am. Chem. Soc., 119, 6384–6391. Copyright (1997) American Chemical Society.)

Equations 2.10 and 2.47 that the difference in peak potentials ΔEp for (i) and (ii), and (i) is given by Equation 2.48 [111]. The relation between Ep − Eo and Kiii for the dimerization reaction (iii) is more complex owing to the presence of the second-order term [B]2 in the expression for Kiii. The reader is referred to the original literature [112] for details, which also contains a zone diagram for this mechanism. Ep − E o = −1.11 ∆Ep =

( )

RT RT  + ln 1 + K CX* nF nF 

( )

RT  ln 1 + K CX* nF 

p

 

p

 

(2.47)

(2.48)

Reversible reactions of type (ii) in Scheme 2.4 are typically observed for radical anions with X being a cation (formation of ion pairs) [114] or a hydroxylic compound such as water or an alcohol

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Techniques for Studies of Electrochemical Reactions in Solution A + ne– + mH+

SCHEME 2.5

B

Electron transfer reaction including reversible proton transfer.

(formation of hydrogen bond complexes) [113,114]. An example of the latter is shown in Figure 2.14b, where also the (larger) effect on the second redox couple is shown. Usually, the product K (CX* ) p is much larger than unity, and consequently a plot of Ep versus logCX* should yield a straight line from the slope of which the stoichiometric number p may be obtained. Deviations from the linear dependence of Ep on logCX* are observed in some cases at high concentrations of X, which may be attributed to the formation of higher associates. Attention should also be brought to a series of papers in which reduction processes including reversible proton transfer reactions (Scheme 2.5) are thoroughly discussed [115] (see also Chapter 13). Reversible dimerizations are observed less frequently, since often the reactions are not fast enough to be treated as thermodynamic equilibria. Examples are the dimerizations of the radical cations of thianthrenes [116] and thiophene derivatives [117]. Let us now reexamine the dimerization mechanisms shown in Scheme 2.3 in a little more detail. The rate-determining steps, (ii) and (iii), in that scheme are both formulated as being irreversible. This, however, is not meant to imply that the dimerization of B, or the coupling of A and B, is an irreversible process by nature. In fact, the only well-documented example of an inherently irreversible reaction of this type is (ii) when B is a neutral free radical. More often, for instance, when B is a radical ion, the chemical reactions (ii) and (iii) in which the intermediates C and I are produced are reversible, and the observed irreversibility is a kinetic phenomenon caused by the further reactions of C and I. Thus, a more natural starting point for the discussion would be the general situation in which both the forward and the backward reactions are considered for all steps except the last one. This is shown in Scheme 2.6. It is convenient first to examine briefly the RS-dimerization mechanism. Usually, the intermediate I is more easy to reduce than A (or to oxidize, when B is a radical cation), which means that the equilibrium constant Kiv for the eCeh step (iv) is large. Also, the rate constant for this type of process is usually large. This allows for the application of the steady-state approximation for the

2B A+B B+I C+X

SCHEME 2.6

(i)

B

A + e– kii k–ii kiii k–iii kiv k–iv kv

C I A+C

(Kii = kii/k–ii)

(ii)

(Kiii = kiii/k–iii)

(iii)

(Kiv = kiv/k–iv)

(iv) (v)

D

Mechanism

Assumption

RR

kv[X] >> k–ii

–d[B]/dt = kobs[B]2

2kii

(vi)

RR

k–ii >> kv[X]

–d[B]/dt = kobs[B]2[X]

2Kiikv

(vii)

RS

kiv[B] >> k–iii

–d[B]/dt = kobs[A][B]

2kiii

(viii)

RS

k–iii >> kiv[B]

–d[B]/dt = kobs[A][B]2

2Kiiikiv

(ix)

Observed rate law

Mechanisms including reversible dimerization.

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kobs

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intermediate I, that is, it is assumed that d[I]/dt = 0. The general rate law resulting for B degenerates to either of two limiting cases depending on the relative magnitudes of k−iii and kiv [B]. When kiv [B] ≫ k−iii, the intermediate I reacts almost exclusively by reaction (iv), and forward reaction (iii) becomes rate determining. The resulting rate law (viii) is the same as rate law (vii) in Scheme 2.3. On the other hand, when k−iii ≫ kiv [B], reaction (iii) behaves as a true equilibrium prior to the ratedetermining reaction (iv) resulting in rate law (ix). It is seen that under conditions where rate law (viii) applies, the reaction order RA/B is equal to 2, whereas rate law (ix) corresponds to RA/B equal to 3. Thus, ideally, it would be possible to distinguish between the two limiting cases by a DCV reaction order analysis carried out as described earlier. Next, let us examine the situation where the reversible RR dimerization (ii) results in a dimer C that is stable on the time scale of the voltammetric experiment. The voltammograms that ideally may be obtained in different kinetic regimes are shown in Figure 2.15a. At low sweep rates (full line), reaction (ii) behaves like a thermodynamic equilibrium, and the response observed 1.0

iν–1/2 (mA V–1/2 s1/2)

0.8 0.6 0.4 0.2 0.0 –0.2 –0.4 –0.6 –0.8 –1.0

–1.2

(a)

–1.4

–1.6

–1.8

–2.0

E (V)

1.1 1.0 0.9 0.5

ipc/ipa

0.8

1.0

0.7

2.0

0.6

5.0

7.0

0.5 0.4 100 400 300 200

0.3 0.2 –6

(b)

–4

10 15 20 40 50

–2 0 –log (RT/nF kf c/v)

2

4

FIgURE 2.15 (a) Simulated (DigiSim®) cyclic voltammograms for a reversible one-electron reduction followed by a reversible RR-dimerization reaction (Scheme 2.5) with Eo = −1.5 V, C A* = 1 mM, Kii = 104 M−1, kii = 106 M−1 s−1, and ν = 0.1 (full line), 30 (dash), and 10,000 (dot) V s−1. For other simulation parameters, see Figure 2.3. Note that units of iν−1/2 are used at the y-axis to facilitate the comparison of voltammograms obtained at different sweep rates. (b) Working curves for the reversible RR-dimerization mechanism (full line) and experimental values (■) of −ipred /ipox for the dimerization of the all-trans 1,10-diphenyl3,8-dimethyldecapentaene radical cations and related compounds. The numbers on the curves refer to the value of K iiC A* (see Scheme 2.5). (With kind permission from Springer Science+Business Media: J. Solid State Electrochem., 2, 1998, 102, Heinze, J., Tschuncky, P., and Smie, A.)

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is close to that shown in Figure 2.14a. At intermediate sweep rates, the dissociation of C is too slow to allow for the observation of oxidation current for B, and at high sweep rates, it is possible even to outrun forward reaction (ii), and the response is essentially that for a reversible electron transfer process. For radical cations, this situation is typically observed when deprotonation of the dimer dication is  slow, and for radical anions under conditions that are free from electrophiles, for example, acids that otherwise would react with the dimer dianion. Most often, this type of process has been observed for radical anions derived from aromatic hydrocarbons carrying a substituent that is strongly electron withdrawing, most notably and well documented for 9-substituted anthracenes [119,120] (see also Chapter 17). Examples from the radical cation chemistry include the dimerization of the 1,5-dithiacyclooctane radical cations [121] and of the radical cations derived from a number of conjugated polyenes [118,122]. The kinetics (kii and k−ii) and thermodynamics (Kii) of such reactions are conveniently studied by CV or DCV. Working curves for CV are shown in Figure 2.15b for different values of Kii together with the experimental points obtained for the oxidation of conjugated polyenes, such as the alltrans 1,10-diphenyl-3,8-dimethyldecapentaene [118]. In going from right to left, corresponding to increasing values of λ, it is seen that the value of −ipred /ipox initially decreases until a minimum is reached. In this kinetic region, the reaction is essentially controlled by the rate of the dimerization. The rising part of the working curve observed after the minimum illustrates the increasing kinetic importance of the dissociation of C to B. In the limit λ = ∞, corresponding to an infinitely large value of kii or concentration and/or an infinitely small value of ν, the chemical reaction responds as a true equilibrium. We now introduce the situation where the dimer C undergoes a subsequent chemical reaction (v). The corresponding working curves, in this case for DCV, are shown in Figure 2.16 for two values of K iiCA* (1 and 10) and different values of kvCX* /k− ii (shown in the figure). The detailed discussion [123] of the curves is beyond the scope of this chapter, but it is seen that the effect of increasing values of kvCX* /k− ii, as intuitively expected, is a gradual transition from the reversible case (kvCX* /k− ii = 0, labeled rev on Figures 2.16a and b) to the totally irreversible case (kvCX* /k− ii = ∞, labeled irr on Figures 2.16a and b). In other words, the reaction of C gradually changes from dissociation to irreversible conversion to D. The latter situation corresponds to the limiting case in which the dimerization is kinetically irreversible as described by rate law (vi) in Scheme 2.6 (=rate law (vi) in Scheme 2.3). 1.0

1.0 K5C(A)° = 1

K5C(A)° = 10

rev 0.8

0.6

0.6

rev

Ri

Ri

0.8

irr

0.4

1

0.1

0.01

0.4

0.01 irr

0.2 0.0 –3 (a)

1

0.1

0.2

–2

0 1 –1 log(k5C(A)°/a)

2

0.0 –3

3 (b)

–2

0 1 –1 log(k5C(A)°/a)

2

3

FIgURE 2.16 DCV working curves for a reversible dimerization followed by an irreversible reaction for (a) K iiC A* = 1 and (b) K iiC A* = 10 (K iiC A* = K5C(A)o in the figures). The labels 1, 0.1, and 0.01 at the working curves refer to the values of kvCX* /k− ii. The rate constant k5 in the figure corresponds to kii in Scheme 2.5, and the constant a is equal to νnF/(RT). (From Hammerich, O. and Nielsen, M.F., Acta Chem. Scand., 52, 831, 1998. With permission.)

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An important consequence of the different shapes of the working curves in Figure 2.16 is that a reaction order analysis based on Equation 2.34 may lead to a value of R A/B that is not related to rate law (2.33) and, hence, cannot easily be used for mechanism assignment. An example serves to illustrate this point. Let it be assumed that the reaction being studied is described by K iiCA* = 1 and kvCX* /k− ii = 0.1 corresponding to the working curve labeled 0.1 in Figure 2.16a. It is seen that RI′ = 0.5 in this case requires that log[ kiiCA* RT /(νnF )] is equal to approximately 0.2. If now the concentration of substrate is increased by a factor of 10, the appropriate working curve is the one labeled 0.1 in Figure 2.16b, and it is seen that RI′ = 0.5 now requires that log[ kiiCA* RT /(νnF )] is equal to approximately −0.8. The two-point value of d log ν1/ 2 /d log CA* calculated from these data is 2.0, which gives R A/B equal to 3.0 using Equation 2.34. This would seem to be in good agreement with rate law (ix) for the RS-mechanism with the electron transfer (iv) being rate determining, but not with the RR-mechanism with irreversible dimerization (ii), for which RA/B is equal to 2. The major origin of the problem is of course that RI′ = 0.5 at the low substrate concentration (Figure 2.16a) corresponds to the part of the working curve that is essentially defined by rate law (vii), whereas RI′ = 0.5 at the high concentration (Figure 2.16b) corresponds to the part essentially defined by rate law (vi). The same type of analysis carried out for kvCX* /k− ii = 0.01 and 1 results in R A/B = 4.1 and 2.2, respectively. It is seen that R A/B increases with decreasing values of kvCX* /k− ii and that therefore virtually any value larger than 2 may result when the analysis is based on Equation 2.34. Problems of this kind may be difficult to detect experimentally. For example, for the reaction following the working curves 0.1, it is seen that the shape of the working curve at low concentration (Figure 2.16a) is similar to that for irr as long as RI′ is lower than 0.7, and at the high concentration (Figure 2.16b), the same is true as long as RI′ is higher than 0.4. Thus, if the data for RI′ recorded during a preliminary DCV investigation were all between 0.4 and 0.7, the results would seem to indicate that the kinetics would be adequately described by a simple rate law such as that in Equation 2.33. The only solution to problems of this kind is to record the full working curves in all cases and to supply the data obtained by DCV (or DPSCA, see Section IV.A) with data from LSV. An example of the latter approach is the study of the dimerization of the radical cations derived from a series of 2,5-diaryl-1,4-dithiins [124]. As a general rule, it is recommended to use as many individual data as possible for mechanistic analyses of this kind, recorded at widely different conditions (concentration, sweep rates, temperature). 6. Irreversible Follow-Up Reactions, the Prepeak Method Reactions between B and X may conveniently be studied also under conditions where CX* /CA* < 1. In that case, the reagent X is consumed during the early parts of the voltammetric wave giving rise to a prepeak (at Ep,pre) in front of the main peak (at Ep,main) for the reduction of A, now in a solution that close to the electrode is essentially free from X [119]. An example of such a voltammogram is shown in Figure 2.17a. The rate constant may be obtained from the potential difference Ep,pre − Ep,main, which has been tabulated for a number of reaction mechanisms [126]. This method has been used extensively in kinetic studies of reactions between radical cations and nucleophiles [126–129]. Alternatively, the determination of k may be based on measurements of Ep/4 − Ep,main, where the quarter-peak potential Ep/4 is the value of E at i = ip,main/4 [125] (see Figure 2.17a). The advantage of this approach is that measurements can be made even in cases where the prepeak appears only as a shoulder on the main peak and where Ep,pre therefore cannot be determined. The working curve for the eC mechanism is shown in Figure 2.17b together with experimental data obtained for the protonation of the anthracene radical anion by benzoic acid. The rate constant resulting from these data is 2.7·106 M−1 s−1 [125]. The prepeak method may be used for reactions with second-order rate constants ranging from approximately 105 M−1 s−1 to that for a diffusion-controlled reaction.

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Ep,main Ep,pre

0.6

i (mA)

0.4

Ep/4 ip,main

0.2

ip,main/4

0.0 –0.2 –0.4 –1.0

–1.2

–1.4

(a)

–1.6

–1.8

–2.0

E (V) log (ν/Vs–1) –2.0

–1.0

0.0

1.0

2.0

3.0

(Ep/4 – Ep) (mV)

150

125

100

75 0.0 (b)

1.0

2.0

3.0

4.0

5.0

log(k4C oHB RT/(vnF ))

FIgURE 2.17 (a) Simulated (DigiSim®) cyclic voltammogram at ν = 1 V s−1 for an eC mechanism with Eo = −1.5 V, C A* = 1 mM, CX* /C A* = 0.5, and k = 109 M−1 s−1. For other simulation parameters, see Figure 2.3. (b) Working curve for the quarter-peak width Ep/4 − Ep,main (bottom scale) together with experimental points (top scale) for the * = 1 mM) in DMSO (rate constant k4). protonation of anthracene radical anion (C A* = 2 mM) with benzoic acid (CHB (From Nielsen, M.F. and Hammerich, O., Acta Chem. Scand., B41, 668, 1987. With permission.)

7. Irreversible Follow-Up Reactions, Redox Catalysis Common to all the methods discussed earlier is that B is generated at the electrode surface, that is, by a direct electron exchange between the electrode and the substrate A. This approach is, however, sometimes hampered by the limitations imposed by the heterogeneous nature of the electron transfer reaction. For instance, studies of the kinetics of fast follow-up reactions may be difficult or even impossible owing to interference from the rate of the heterogeneous electron transfer process. In such cases, the kinetics of the follow-up reactions may be studied instead by an indirect method, generally known as redox catalysis [2,130–132]. Another application of redox catalysis is the study of dissociative electron transfer reactions (Scheme 2.1, see also Chapter 14) [80,81,133,134]. The principle is illustrated in Scheme 2.7 for a reaction that in a direct process follows the eC mechanism. The direct process includes the reduction of A to B (i) followed by the irreversible reaction of B to a product C (iv). If now a compound P, a so-called mediator, that is easier to reduce than A is

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Organic Electrochemistry A + e–

B

E oAB

(i)

P + e–

Q

E oPQ

(ii)

P+B

Kiii

(iii)

kiii

Q+A B

SCHEME 2.7

kiv

k–iii

(iv)

C

Mechanism including redox catalysis.

added to the solution, the direct reduction of A is replaced by the reaction sequence (ii) and (iii). During the electron transfer process (iii), the oxidized form P of the mediator is regenerated. In other words, the reduction of A to B is catalyzed by the redox couple P/Q. Cyclic voltammograms typical for the catalyzed process are shown in Figure 2.18. The effect of increasing values of kiii is shown in Figure 2.18a, and the effect of an increasing concentration ratio CA* /CP* is shown in Figure 2.18b.

4

i (mA)

3

ip,revγ

2

ip,cat

1

ip,rev 0 –1 –1.0 (a)

–1.2

–1.4 E (V)

–1.6

–1.8

25 20

i (mA)

15 10 5 0 –5 –1.2 (b)

–1.3

–1.4

–1.5 E (V)

–1.6

–1.7

–1.8

FIgURE 2.18 Simulated (DigiSim®) cyclic voltammograms at ν = 1 V s−1 for the catalytic reaction shown o in Scheme 2.7 with EPQ = −1.5 V, CP* = 1 mM, Kiii = 5·10 −7, and (a) k iv = 108 M−1 s−1, C A* /CP* = 4 , and k iii = 0 4 5 (full line), 10 (dash), 10 (dot), 10 6 (dash-dot), and 108 (dash-dot-dot) M−1 s−1, and (b) k iii = 10 6 M−1 s−1, k iv = 108 M−1 s−1, and C A* /CP* = 1 (full line), 2 (dash), 5 (dot), 10 (dash-dot), 20 (dash-dot-dot), and 50 (short dash). For other simulation parameters, see Figure 2.3. The magnitude of ip,cat is shown for k iii = 105 M−1 s−1.

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The description of the kinetics is simplified considerably by the fact that often the three rate constants, kiii, k−iii, and kiv, are all large allowing for the application of the steady-state approximation for the concentration of B, that is, −d[B]/dt = 0 [124]. The system is conveniently discussed by the introduction of three dimensionless parameters: (1) the kinetic parameter λ, Equation 2.49; (2) the competition parameter σ, Equation 2.50; and (3) the concentration excess factor γ, Equation 2.51. In addition, it is convenient to introduce the catalytic efficiency CAT, Equation 2.52, where ip,cat is the peak current during LSV observed for the catalytic system and ip,rev is the peak current observed for P in the absence of A (see Figure 2.18a). The product ip,revγ then is the maximum possible catalytic current. λ iii =

σ=

kiiiCP* RT νF

(2.49)

kiv k− iiiCP*

(2.50)

CA* CP*

(2.51)

γ=

CAT =

ip,cat ip,rev γ

(2.52)

Two limiting cases result depending on the magnitude of σ. For σ ≫ 1, the kinetics are controlled by kiii, that is, by the rate of the solution electron transfer reaction, and for σ ≪ 1 by Kiiikiv, that is, the case where the electron transfer reaction may be considered as a true equilibrium prior to the reaction of B. The transition is shown by the voltammograms in Figure 2.18a for σ varying between 5·102 (lower curve for kiii = 102 M−1 s−1) and 5·10 –4 (upper curve for kiii = 108 M−1 s−1). Under conditions where σ ≫ 1, the value of kiii may be obtained by fitting the appropriate working curve for a certain CAT(λiii, γ) [130] to the experimental data. In this kinetic region, the value of CAT is proportional with CP* . On the other hand, for σ ≪ 1, the value of Kiiikiv may be obtained by fitting the working curve for CAT(λiii, σ, γ) [130]. In this region, the value of CAT is independent of CP* . For the complete characterization of the system, the value of k−iii is needed. This may be obtained under conditions where the system is under mixed control of reactions (iii) and (iv), again by fitting the appropriate working curve to the experimental data. Once Kiii is known, it is possible also to o determine EAB as follows: o o EAB = EPQ +

RT ln K iii nF

(2.53)

Numerous examples of the application of redox catalysis to studies of the kinetics of the cleavage of the carbon–halogen bond in the radical anions of aromatic halides have been reported [131,135] (see also Chapter 24). Other examples include the fragmentation of tert-butyl radicals from the radical cations of NADH model compounds [136], the reduction of CO2 [137], and the cleavage of the carbon–nitrogen bond in the radical anions of 1,1-dinitroalkanes [138]. Another application of redox catalysis is the study of dissociative electron transfer reactions (Scheme 2.8). The resulting free radical R· may undergo either of two reactions, coupling with the mediator radical anion (iii) or reduction to R− (iv) [134]. The coupling reaction is usually considered as unwanted since the mediator is irreversibly consumed in this step. The reaction is, however, synthetically useful [134].

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Organic Electrochemistry A + e– A –+ R – X

A

A – (EoA/A –) kii

(i)

A + R + X–

(ii)

kiii

A – R–

(iii)

kiv

A + R–

(iv)

–+ R

SCHEME 2.8 Reduction of alkyl halides by redox catalysis.

The competition between reactions (iii) and (iv) is usually described by the magnitude of the parameter q given as follows, which may be determined, for instance, by LSV [139,140]: q=

kiv kiii + kiv

(2.54)

It is seen that q = 0 for a clean coupling reaction and that q = 1 for a clean electron transfer reaction. The value of kiii is close to that for a diffusion-controlled process, whereas the value of kiv increases o when EA/ moves in the negative direction. Accordingly, a plot of q versus EAo /A• − will give rise A•− q corresponding to q = 0.5 may be determined. to an S-shaped curve from which the potential E1/2 q Once the value of E1/2 is known, the standard potential for the reduction of R· may be determined by application of the Marcus theory [134]. 8.

Chemical Reactions Following Quasireversible or Irreversible Electron Transfer Reactions Although processes of this kind are often met in practical work, the presentation of the theoretical data associated with these reactions is beyond the scope of this chapter. The reader is referred to the original literature for details [19,141]. When studied under purely kinetic conditions as defined earlier, processes belonging to this category are usually observed to result in values of Ep − Ep/2 larger than those given, for instance, in Equations 2.43 and 2.44. The analysis of the experimental data typically includes the simulation of working curves that take into account also the effect of ko and α. In spite of the fact that the description of the system now involves ko and α in addition to the usual kinetic parameters, it is often possible to gain information about all the parameters by a strategic variation of ν, CA* , and CX* . An illustrative example is the dimerization of 2,5-diaryl-1,4-dithiin radical cations [124]. 9. general Remarks and Conclusions The applications of LSV and CV to the study of chemical processes following an electron transfer reaction are so numerous that a review of the subject is clearly beyond the scope of this chapter. The examples were selected to demonstrate the application of the techniques in practical work. Although obvious, it should be emphasized that electrochemical reactions are not different from any other chemical reaction and, therefore, that the whole arsenal of methods of attack known from conventional kinetics may be used in the characterization of the process. This includes, not least, the effect of temperature [100,127,142–144] and studies of kinetic isotope effects [144,145].

E.

LIMITING EXPERIMENTAL FACTORS

The shape and position of the voltammogram depends not only on the kinetics and thermodynamics of the electrochemical reaction, but is affected also by the design of the voltammetry cell and the potentiostat. Closely related to this is the question of measurement precision.

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1. The Cell and the Working Electrode Two problems may arise at low ν. The first of these is related to the assumption that mass transport is caused by diffusion only. If we approach, during a voltage sweep, potentials that cause conversion of A, a diffusion layer starts to develop. At a planar electrode, assuming linear diffusion, its thickness δ increases with time t as given in the following: δ = πDAt

(2.55)

This Nernstian definition of δ assumes a linear approximation of the concentration profile with its slope equal to that found at the electrode surface [146]. Other formulations are possible [146], and one of them is used in Section IV.A.3 in Chapter 1, Equation 1.149. When δ becomes larger than a few tenths of a millimeter, natural convection begins to interfere, and the assumption of diffusion as the only means of mass transport is no longer valid. For typical values of DA (10 –5 cm2 s−1), this happens when t is of the order of 10 s, corresponding to ν = 0.2 V s−1 for a CV sweep with Einitial − Eswitch = 1 V. At times larger than approximately 1 min, or ν approximately less than 0.025 V s−1, the deviations from pure diffusion are so pronounced and unpredictable that the measured current cannot be related to a practical theoretical model. Only recently, the effect of natural convection on CV was modeled by a stagnant convection layer [147]. This influence of natural convection at low ν may be reduced by using properly shielded working electrodes [148] (Figure 2.19), which are, however, difficult to handle in practical work. In particular, the replacement of electrolyte in the diffusion layer after recording a cyclic voltammogram proves impractical. Another problem at low ν is the so-called edge-effect observed when the thickness of the diffusion layer at an unshielded working electrode is comparable to its diameter [149,150]. At the edges, the diffusion of material to and from the electrode is not limited to being linear as illustrated in Figure 2.20, and it is seen that the deviations become progressively larger when the diameter of the electrode becomes smaller until, in the limit, the transport may be described as hemispherical diffusion. This applies for the so-called ultramicroelectrodes, and it is advantageously used with such devices to improve diffusional transport to the electroactive surface (see Section III). By inspection of Figure 2.20, it is easily understood that mass transport close to the edges is more effective than at the central part of a circular electrode, and accordingly, the current density (e.g., in units of mA cm−2) increases when the diameter of the electrode decreases. The effect of this change in the diffusion pattern on the voltammogram for a reversible electron transfer reaction, Equation 2.1, is shown in Figure 2.21. It is seen that the gradual change of the diffusion pattern has a profound influence not only on the magnitude of the current density but also on the appearance of the voltammogram, which gradually changes to the S-shaped curve characteristic

FIgURE 2.19 Schematic of a shielded working electrode. The arrows illustrate the diffusion pattern below the electrode.

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FIgURE 2.20 Schematic of the diffusion pattern below electrodes with decreasing diameters of the electrode surface (from left to right).

for hemispherical diffusion [150]. In the latter case, the diffusion of material away from the electrode is so effective that a reverse current for the reoxidation of B to A cannot be observed even, as in this case, when B is perfectly stable in the solution. At high ν, the relative positions of the three electrodes need to be optimized and the size of the working electrode to be reduced in order to minimize resistance and capacitance problems [151]. The quality of the potentiostat, usually reflected by the rise time, also becomes critical at high ν. The two factors are actually not separable since the electrochemical cell is inherently a part of the total electronic circuit. 2. Uncompensated Solution Resistance, Ru The Ru problem in LSV and CV has received considerable attention [152–159]. The electrode potential during LSV for a sweep in the negative direction is described by Equation 2.56, where i is the current flowing at the time t. Emeasured = Einitial − νt + iRu

(2.56)

This equation suggests that the real electrode potential Ereal can be determined only under conditions where the last term (the iR-drop) vanishes. A first countermeasure is to reduce Ru by minimizing the distance between working electrode and reference electrode tip, while avoiding diffusion shielding and increasing the conductivity of the electrolyte. Then, the concentration of the electroactive species may be decreased, to decrease i. In many situations, especially in organic electrolytes, still an appreciable iRu remains. Thus, the problem is to correct Emeasured for the contribution of iRu. This is normally accomplished by a positive feedback circuit incorporated into the potentiostat, which adds a fraction of the current follower output to the voltage provided by the function generator. If the feedback resistance Rf is exactly equal to Ru, the iRu term in Equation 2.56 is compensated for and Emeasured is equal to Ereal. The problem then is the selection of the value of Rf. Although this can be accomplished by direct measurement of Ru and other techniques [152,153], a simpler procedure is desirable for the level of sophistication of most electrochemical studies. Such a simple and convenient method is to adjust the feedback circuit until the output of the current follower goes into oscillation and then to back off slightly [159]. This method is quite effective for electrodes with diameters less than 2 mm as long as ν is not larger than about 100 V s−1. For high sweep rate experiments, specialized iRu-drop correction techniques have been employed [49b,d,e]. Unfortunately, the iRu effect in CV affects the shape of the current/potential curves in a way similar to the effect of quasireversibility (increase in ΔEp; see Section II.D.8). How, then, can we prove the absence of iRu for quasireversible voltammograms? Since iRu increases for a given electrolyte and a fixed positioning of the electrodes with i, any such artifact linearly increases with the substrate concentration C*A. Within a series of voltammograms

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Techniques for Studies of Electrochemical Reactions in Solution 0.8

iA–1 (mA cm–2)

0.6 0.4 0.2 0.0 –0.2 –0.4 –0.6 –1.0

–1.2

–1.4

–1.6

–1.8

–2.0

–1.6

–1.8

–2.0

–1.6

–1.8

–2.0

–1.6

–1.8

–2.0

E (V)

(a)

1.0

iA–1 (mA cm–2)

0.8 0.6 0.4 0.2 0.0 –0.2 –0.4 –0.6 –1.0

–1.2

–1.4

(b)

E (V)

3.0

iA–1 (mA cm–2)

2.5 2.0 1.5 1.0 0.5 0.0 –0.5 –1.0 –1.0

–1.2

–1.4 E (V)

(c) 24

iA–1 (mA cm–2)

20 16 12 8 4 0 –4 –1.0 (d)

–1.2

–1.4

E (V)

FIgURE 2.21 Simulated (DigiSim®) cyclic voltammograms for a reversible one-electron reduction with Eo = −1.5 V at electrodes with (a) r = 0.7 mm, (b) r = 0.05 mm, (c) r = 0.01 mm, (d) r = 0.001 mm. Note that the y-axis displays the current density (in mA cm−2) in order to facilitate the comparison of voltammograms obtained at different surface radii. For other simulation parameters, see Figure 2.3.

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recorded at different C*A , ΔEp should be constant if any iRu drop is absent. On the other hand, a concentration dependence of ΔEp indicates that iRu significantly contributes to the experimental outcome. Note that this approach is no longer valid if reactions with second- or higher-order kinetics are coupled to the electron transfer. The concentration dependence of iRu has further been used to correct ΔEp values recorded as a function of v and C*A [160]. 3. Importance of Precision in Potential Measurements The precision of Ep required for comparison with literature values is of the order of ±5 mV. This can easily be accomplished by commercial computer software. In other cases, including LSV analysis (see Section II.D.4) and studies of the effect of small structural changes such as deuterium isotope effects, the required precision is ±0.5–1 mV or better. This can easily be appreciated by considering that the theoretical values of, for instance, dEp/dlogν in LSV, Equation 2.40, differ by only 10 mV for the first- and second-order follow-up reactions. Similarly, it has been found that 2,7-d2-pyrene is only 1.5 ± 0.2 mV more difficult to reduce than pyrene itself [161]. Such a high precision requires highly stable reference electrodes and advanced data treatment including averaging of measurements, baseline corrections, and sophisticated peak fitting routines. Another approach involves the application of derivative LSV or DCV (Figure 2.10b). These techniques allow a precision of the order of ±0.2 mV to be realized during peak potential measurements [28–30,35]. The feature giving the high precision is that the peak potential is measured at the point where the steeply descending first derivative crosses zero. 4. background Currents The faradaic current, which is usually of interest in the analysis of cyclic voltammograms with respect to electron transfer–coupled organic reactions, is only one contribution of the total current through the electrode. Other components are related to double-layer charging [158] and impurities. Conventionally, these unwanted currents are designated as background current. For precise current data, as well as an adequate analysis of the shape of a cyclic voltammogram, the background current must be subtracted from the raw current/potential curve. Most commercial instruments have simple procedures for this. However, the respective background curve must be recorded under identical or at least as similar as possible conditions as the substrate curve. This can be achieved by recording the background first and then adding the substrate to the electrolyte in the cell (possibly in several portions to generate solutions of increasing concentration after each addition) without changing the cell geometry. Still, particularly for signals close to the limits of the accessible potential window (low signal-to-background ratio), artifacts may be introduced by subtraction. Detrimental effects can also be observed if the substrate slightly changes the double layer, for example, by adsorption phenomena.

F. COMPUTER-BASED METHODS FOR ANALYSIS OF VOLTAMMETRIC DATA Although relevant information can be extracted from voltammograms by the techniques described in Sections II.C and II.D already, the analysis of CV data is traditionally linked to computer application. This supports data acquisition, signal processing, and the determination of kinetic and thermodynamic parameters of interest. It is particularly true for cases where the relation between data and information is nonlinear, the chemistry is described by several parameters (e.g., rate constants), and these are possibly correlated. Then, the task at hand (sometimes called the inverse problem) may need extensive computation. Some aspects with respect to simulation of electrochemical processes are discussed in Section III.D in Chapter 5. Here, we restrict the treatment to practical comments and the case of CV.

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Several approaches were already developed several decades ago and include, for instance, pattern recognition [162], the recording of current–time profiles of the form Δlni/Δlnt versus t for mechanistic classification [163], as well as nonlinear regression techniques [164–167]. Efforts have also been made to use knowledge-based systems for the elucidation of reaction mechanisms [168]. More recently, artificial neural networks [169] and wavelet transformation [170] were applied, but often, such work is devoted mainly to analytical electrochemistry. The application of support vector machines showed that analysis of full voltammograms rather than only selected prominent points (peaks) provides better estimation of diffusion coefficients [171]. As examples for more widely used advanced techniques, we discuss convolution and fitting approaches. 1. Semi-Integration and Convolution Techniques The essential feature of these methods is that semi-integration [172–177] or calculation of the convolution integrals, Equation 2.57 [156,178–184], results in the transformation of the voltammogram into what has been called a neopolarogram. Experimentally, the convolution integral may be evaluated by using the Riemann–Liouville algorithm, Equation 2.58 (see also Reference 1). The resulting curve is given by Equation 2.59, where Il is defined in Equation 2.60. t

1 I (t ) = 1/ 2 π

i(u)

∫ (t − u )

1/ 2

(2.57)

du

0

1/ 2 t / ∆

∆ I (t ) =   π

1/ 2  t  t (i( j∆) + i( j∆ − ∆))   − j + 1  −  −  ∆  ∆ j =1 

∑

E = E1/2 +

RT I l − I (t ) ln nF I (t )

I l = nFADA1/ 2CA*

1/ 2

 j 

   

(2.58)

(2.59)

(2.60)

Logarithmic analyses can be carried out on the neopolarogram in much the same way as with classical polarography (see Section V). Examples that illustrate the application of the convolution technique include not only simple quasireversible electron transfers [185] but also more complex cases such as the reversible dimerization of radical cations [121], the study of dissociative electron transfer reactions [186,187], and investigations of the possible potential dependence of the transfer coefficient α [184,188,189]. 2. Fitting Simulated to Experimental Voltammograms The advances in computer technology have, in more general terms, stimulated the development of techniques that use more than just a few data points, such as Ep, Ep/2, and ip, along the voltammetric wave. Thus, after selection of a proper mechanistic hypothesis, a voltammogram is calculated (predicted; see Chapter 5) and, by means of variation of the kinetic, thermodynamic, and transport parameters provided to the simulation, fitted to the respective experimental curve. This would at first glance seem to be an advantage. However, it should be brought in mind that the problems originating from, for instance, background currents and nonproper adjustments of the electronic equipment remain and it becomes even more crucial to avoid these artifacts. For less ideal experimental data, some skill is required to incorporate parameters, for instance, for the uncompensated electrical solution resistance (see Section II.E.2) and the double-layer capacity (see Chapter 1) in the data treatment. Furthermore, this approach can be dangerous if too many unknown parameters are

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35 30 25

i (μA)

20 15 10 5 0 –5 0.0

0.2

0.4 E (V vs.

0.6

0.8

1.0

Ag/Ag+)

FIgURE 2.22 Cyclic voltammetry data (solid line) for the oxidation of 2,4-dimethyl-3-ethylpyrrole in MeCN at ν = 0.2 V s−1 and the simulated curve corresponding to the mechanism shown in Scheme 2.9. (Reprinted from Hansen, G.H. et al., Electrochim. Acta, 50, 4936, 2005.)

subject to (possibly automated) variation, and overfitting may pretend good coincidence between experimental and fitted curves while the parameter values are not reliable (see also recommendations in Section III.D in Chapter 5). The opinion of the authors of this chapter is that trends in the kinetic data and their origin are often easier to locate when a more classical approach to the data analysis is taken. Still, with appropriate care, fitting voltammograms can be a valuable tool. It should be mentioned here also that the commercial software package DigiSim® offers the possibility to fit simulated voltammograms directly to experimental ones. A particular problem is that voltammograms, even for very different mechanisms, are very similar, and for that reason, the method should also be used with care. Also in this case, it is recommended that voltammograms obtained with different scan rates and a variety of substrate and reagent concentrations are included in the fitting procedure. In our opinion, the method works best when the voltammogram presents structural features that are unique for the mechanism in question. An example, the oxidation of 2,4-dimethyl-3-ethylpyrrole, is shown in Figure 2.22. The additional structural feature in this case is the trace crossing observed at 0.6 and 0.15 V, respectively, which is caused by slow proton transfer from the initially formed dimer to the unoxidized substrate. The overall reaction scheme is summarized in Scheme 2.9. The values of the rate and equilibrium constants for the proton transfer reactions resulting from the curve fitting were k2+,m = 6∙103 M−1 s−1, K2+,m = 7 and k+,m = 4∙103 M−1 s−1, K+,m = 0.15, and it is seen that k2+,m and K2+,m are both found to be larger than k+,m and K+,m, respectively, as intuitively expected. The values of the two proton transfer rate constants compare favorably with those already reported for alkylpyrroles [191,192].

III.

ULTRAMICROELECTRODES AND SCANNINg ELECTROCHEMICAL MICROSCOPy

A.

ULTRAMICROELECTRODES

The discussion in the Section II is mainly concerned with electroorganic reactions at macroscopic electrodes with sizes in the mm range. Then, linear diffusion with flux of material perpendicular to

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+

2

Me

Et

Me

Et

–2e 2e–

N H PyrH

2 Me

+

2

Me

+N H H Me Et +PyrH-HPyr+

Et Me

H N + Me

Et

Me

+

Me Me

Et

Me

Et

Et Me

Me N H

Et

Me

Me

Me H N

–e– Me

e–

Et

N H

Me

Me Et

Pyr-Pyr

Me

Me H N+ H H

Me Me

Et

PyrH2+

Et –e– e–

Me

Me +N H

H N + Me Me

Et

+Pyr-Pyr+

Pyr-Pyr+

Et

Et +

Me

Pyr-Pyr

PyrH

H N

Me N H

N H

Pyr-HPyr+

H N+ H H

Me

PyrH2+

H N

Me

Me

Et

k+,m and K+,m

+ Me

+

Pyr–HPyr+

Et H N + Me

Et

H N + Me

N H H Me

N H PyrH

Me N H H Me

Me

k2+,m; K2+,m

+

Me

Et

Me H

H N + Me

+N H H Me Et +PyrH-HPyr+

N H PyrH+

Et

Me H

kdim Me

N H PyrH+

Et

Me

Et

Me

–

Me H N+ +N H

Me

kp Me

Products

Et

+Pyr-Pyr+

SCHEME 2.9 The initial steps for the oxidation of 2,4-dimethyl-3-ethylpyrrole (kryptopyrrole). (Adapted from Hansen, G.H. et al., Electrochim. Acta, 50, 4936, 2005.)

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the electrode surface is ensured. If the electrode size becomes smaller (Figure 2.20) and reaches, within the time scale of the experiment, the thickness of the diffusion layer (see Equation 2.55), linear diffusion can no longer be assumed, and the term ultramicroelectrode is used. Some aspects of these devices are covered in Chapters 1 and 10. Nowadays, even electrodes with low nanometric dimensions (nanoelectrodes) are prepared and used in electrochemical studies [193]. The definition of an ultramicroelectrode is somewhat arbitrary and, with reference to the common disk geometry, is currently meant to describe electrodes with diameters of about 10–20 μm or less [131,150,151,194,195]. The electrochemical response at ultramicroelectrodes can be classified into two categories depending on the value of ν, and a rational definition of ultramicroelectrode behavior may be based on these [196]. At low ν, the transport to and from an ultramicroelectrode is best described as hemispherical diffusion; this results in a faradaic current that greatly exceeds that expected for linear diffusion [150,197] (see Section II.E.1 and Figure 2.21). An important feature of the voltammogram shown in Figure 2.21d is the absence of a peak. Instead, the current reaches a plateau indicating that a steady-state has been obtained. The steady-state current for an ultramicroelectrode inserted in a large insulating shaft (Figure 2.20c) is given by Equation 2.61, where r is the radius of the electrode surface [198]. The effective transport resulting from hemispherical diffusion also results in an electrode system that is under certain conditions relatively insensitive to natural convection [196]: i = 4nFrDACA*

(2.61)

The second category of electrode response for an ultramicroelectrode occurs at high ν. Under these conditions, linear diffusion is operative, and the response does not differ from that of conventional electrodes with surface diameters in the 0.1–2 mm range. However, the effects of double-layer charging (see Chapter 1) and solution resistance are considerably reduced owing to the small size of the electrode. The transition from low to high ν is shown in Figure 2.23 for an electrode with r = 1 μm at ν = 0.03, 10, 500, and 100,000 V s−1, respectively. Although the use of ultramicroelectrodes is not restricted to any specific measurement technique [131,151,194,199], only applications in the context of CV at high sweep rates will be considered here (see also Section IV). For the studies of reaction kinetics using ultramicroelectrodes under steady-state conditions, the reader is referred to the original literature [150,195,200] and Section III.B. 1. Determination of Eo′ for the A/b Redox Couple When b Is Highly Reactive The perspective in using very high voltage sweep rates during CV has been demonstrated by several authors [131,151,201–205]. Historically, it is of interest to notice that Perone already in 1966 used sweep rates up to 50,000 V s−1 [201] in studies of the electron transfer rates for inorganic redox couples, the 1,000,000 V s−1 mark was passed in 1988 [204,205], and 10 MV s−1 have been reached recently (see Section II.B.3). Cyclic voltammograms for the reduction of anthracene in MeCN at sweep rates up to 100,000 V s−1 [203] are illustrated in Figure 2.24. It is seen that the contribution from the double-layer charging current increases relative to the contribution from the faradaic current with increasing sweep rate. This is because the double-layer charging current is directly proportional to the sweep rate, whereas the faradaic current increases only with the square root of the sweep rate (Equation 2.13). Subtraction of a background curve obtained in the absence of the substrate may eliminate the problem [206]. An obvious application of high sweep rates is the determination of Eo values for A/B couples where B undergoes a chemical reaction so fast that it cannot be outrun by the sweep rates applicable to conventional electrodes [49a,101,207]. Once Eo is known, the rate constant may be determined from LSV relations of the type already given (Equations 2.45 and 2.46).

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iν–1/2 (nA V–1/2s1/2)

4 3 2 1 0 –1 –1.0

–1.2

–1.4

–1.6

–1.8

–2.0

E (V)

(a)

iν–1/2 (nA V–1/2s1/2)

0.3

0.2

0.1 0.0 –0.1 –1.0

–1.2

–1.4

–1.6 E (V)

–1.8

–2.0

–1.2

–1.4

–1.6 E (V)

–1.8

–2.0

–1.2

–1.4

–1.6

–1.8

–2.0

(b)

0.100

iν–1/2 (nA V–1/2s1/2)

0.075 0.050 0.025 0.000 –0.025 –0.050 –1.0 (c)

iν–1/2 (nA V–1/2 s1/2)

0.075 0.050 0.025 0.000 –0.025 –0.050 –1.0 (d)

E (V)

FIgURE 2.23 Simulated (DigiSim) cyclic voltammograms for a reversible one-electron reduction with E o = −1.5 V at an electrode with r = 0.001 mm and (a) ν = 0.03, (b) ν = 10, (c) ν = 500, (d) ν = 100,000 V s−1. For other simulation parameters, see Figure 2.3. Note that units of iν−1/2 are used at the y-axis to facilitate the comparison of voltammograms obtained at different sweep rates.

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(a–d)

200 nA

–2.5

–2.0 E (V) vs Ag/AgClO4

(e)

(g)

(f )

900 nA

500 nA –2.0

–2.5

–2.0 –2.5 E (V) Ag/AgClO4

1600 nA –2.0

–2.5

FIgURE 2.24 Cyclic voltammograms for the reduction of anthracene (2.22  mM) in MeCN containing Et4NClO4 (0.6 M) at a gold electrode with r = 6.5 μm and (a) ν = 1,000, (b) ν = 2,000, (c) ν = 5,000, (d) ν = 10,000, (e) ν = 20,000, (f) ν = 50,000, and (g) ν = 100,000 V s−1. (Reprinted with permission from Howell, J.O. and Wightman, R.M., Anal. Chem., 56, 524–529. Copyright (1984) American Chemical Society.)

2. Determination of Heterogeneous Electron Transfer Rate Constant It follows from the discussion of electron transfer reactions in Section II.C that a reversible process when studied by CV inevitably passes into the quasireversible regime at some value of ν when ν is allowed to increase. For ko = 3 cm s−1, a value typical for many aromatic hydrocarbons, for example [71], it is seen from Equation 2.27 that this happens at approximately ν = 170 V s−1. Thus, studies of electron transfer rates in this region require voltage sweep rates in the range 200–1000 V s−1. Ultramicroelectrodes are superb for this purpose as demonstrated, for example, in a study of the oxidation of ferrocene [208]. Peak potential separations recorded at sweep rates between 500 and 3000 V s−1 at an electrode with a surface diameter of 10 μm over a range of substrate concentrations resulted in k o equal to 1.10, 1.13, and 1.13 cm s−1 (see, however, a critical review that compares such values determined under various conditions [209]). A more recent example is the oxidation of selenanthrene and related compounds [210]. 3. Kinetics of Rapid Follow-Up Reactions Ultramicroelectrodes are excellent tools also for the study of the kinetics of fast follow-up reactions provided that the high sweep rates needed do not bring the electron transfer process into the quasireversible region. Thus, a reaction that, during a study with conventional electrodes, is under

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Techniques for Studies of Electrochemical Reactions in Solution

139

purely kinetic conditions may at high sweep rates be studied under mixed diffusion and kinetic control using the approach described in Section II.D.3. An early example of this approach is the study of the interconversion of the A and B forms of bianthrone [211]. The stable isomer at ambient temperature is the yellow A form, which is equilibrated with the green B form as the temperature is increased. Reduction of A produces A• −, which rapidly undergoes a conformational change to the B-like radical anion. The formal potentials for the two forms differ by about 230 mV for the isomers of 1,1′-dimethylbianthrone in DMF at 361 K. The rate constant for the conversion of B to A could be determined to be equal to 500 s−1 under these conditions by comparing experimental cyclic voltammograms obtained at 200, 500, 1,000, and 5,000 V s−1, using 5 μm Pt disk electrodes, to theoretical voltammograms obtained by digital simulation. Other examples include studies of the cleavage of the carbon–halogen bond in the radical anions of aromatic halides [205,212] and the RR-dimerizations of the N,N-dimethylaniline radical cation [213] and diphenylamine radical cation [214]. 4. Voltammetry in Electrolytes with Low Conductivity Already early experiments with ultramicroelectrodes proved the usability in situations of extremely low conductivity: thus, electrochemistry in the gas phase was accomplished [215]. The low absolute currents through these devices can be sustained without the use of a supporting electrolyte. In addition, the iRu drop at a disk ultramicroelectrode is proportional to the radius [215b] and very low for small sizes, even if the uncompensated resistance Ru is large. Thus, voltammetry in solutions with low concentrations of supporting electrolytes is based on diffusional and migrational transport and has been used to study the effects of the ionic concentration and of ion pairing on organic electrode reactions [216] and their mechanisms (e.g., aryl halide bond cleavage [216c] or anthraquinone radical anion disproportionation [216b]).

B.

APPLICATION OF SCANNING ELECTROCHEMICAL MICROSCOPY STUDIES OF REACTION KINETICS

FOR

Scanning electrochemical microscopy (SECM) [217] is a member of the growing family of scanning probe techniques. In SECM, the tip serves as an ultramicroelectrode at which, for instance, a radical ion may be generated at very short distances from the counter electrode under steady-state conditions. The use of SECM for the study of the kinetics of chemical reactions following the electron transfer at an electrode [217] involves the SECM in the so-called feedback mode [218], where the counter electrode serves to collect the primary electrode intermediate generated at the ultramicroelectrode (Figure 2.25). At sufficiently small distances between the ultramicroelectrode tip and the counter electrode, the time of diffusion from the tip to the counter electrode is of the same order as the lifetime of the intermediate, which makes possible the detection of the intermediate at the counter electrode. The upper limit for first-order rate constants is close to 105 s−1 and for second-order rate constants 108–109 M−1 s−1 [219]. The relationship between the distance and the kinetics of the follow-up reaction is expressed by the dimensionless parameter K, as follows, where k is the rate constant of the chemical reaction, r is the radius of the tip, and d is the distance between the tip and the counter electrode: K=

K=

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kd 3 rDA

kCA* d 3 rDA

(first-order rate law)

(2.62)

(second-order rate law)

(2.63)

140

Organic Electrochemistry

Tip of ultramicroelectrode

e–

k

A

B

Products

e–

Counter electrode

FIgURE 2.25 The principle of the application of scanning electrochemical microscopy for studies of reaction kinetics.

At very short distances, depending on the rate of the chemical reaction, essentially all B generated at the tip is detected by the counter electrode. When the distance gradually is made larger, an increasing fraction of B reacts to products until, in the limit, no B is detected at the counter electrode. By plotting the ratio between the current at the counter electrode ic and the tip current it, −ic/it, as a function of the distance, information about the kinetics of the follow-up reaction may be obtained by fitting the data to the appropriate working curve −ic/it versus log  K (Figure 2.26). Thus, the parameter −ic/it plays a role equivalent to that, for instance, of the collection efficiency in rotating ring-disk electrode (RRDE) voltammetry (Section VI.C). The distance between the tip and the counter electrode is determined by a calibration procedure including the recording of −ic/it for a reversible redox couple for which B is stable and fitting the data to a working curve [220]. An advantage of SECM is that it is a steady-state technique that does not require the measurements of transient currents. The technique also offers the same advantage as, for instance, double potential-step chronoamperometry (DPSCA) (Section IV.A) in that the potential of the tip electrode can be made sufficiently negative that the heterogeneous electron transfer process proceeds at the diffusion-controlled rate. An example of the application is the oxidation of trans-anethole [trans-1-(4-methoxyphenyl) propene], the radical cation of which by LSV was found to undergo rapid RR-dimerization (through C-2). The rate constant measured by SECM is as high as 4·108 M−1 s−1 [221]. The fit of the experimental data to the working curve is shown in Figure 2.26. The SECM technique has also been used to determine the rate constant for the dimerization of the radical anions of acrylonitrile, the prototype example of electrohydrodimerization (Chapter 17). The value of k was found to be 6·107 M−1 s−1 [222]. SECM-like methodology was used to map concentration profiles in front of a larger electrode [223]. Recently, an extensive compilation of SECM modes and their application to determine reaction rates with lateral resolution across a surface has been provided [224], including those of monolayers, polymers, bioactive materials, membranes and tissues, enzymes, and individual cells.

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141

Techniques for Studies of Electrochemical Reactions in Solution 1 0.9 0.8 0.7

Is/IT

0.6 0.5 0.4 0.3 0.2 0.1 0 –1

–0.5

0

0.5

1

1.5

2

2.5

3

3.5

4

log(K)

FIgURE 2.26 SECM working curve for the RR-dimerization mechanism together with experimental points obtained for the oxidation of trans-anethole. The ratio IS/IT at the figure corresponds to −ic/it in the text, and K is defined in Equation 2.63. (From Demaille, C. and Bard, A.J., Acta Chem. Scand., 53, 842, 1999. With permission.)

IV. POTENTIAL-STEP AND CURRENT-STEP METHODS A.

CHRONOAMPEROMETRY AND DOUBLE POTENTIAL-STEP CHRONOAMPEROMETRY

When the potential of a planar electrode is suddenly shifted from a value at which no current flows to a value where the electron transfer, Equation 2.1, proceeds at the diffusion controlled rate, the current–time behavior is given as follows, where ic is the double-layer charging current [1,152]: i=

nFACA* DA1/ 2 + ic (πt )1/ 2

(2.64)

When conventional electrodes with diameters between 0.1 and 2 mm are used, the latter quantity has usually decayed to zero after 0.5 ms or less and may be neglected in experiments lasting 1 ms or more. This decay time is reduced to the microsecond time regime when ultramicroelectrodes are used [101,131,225]. According to Equation 2.64, which for ic = 0 is known as the Cottrell equation, the current approaches zero when the time approaches infinity. However, undisturbed linear diffusion can be maintained only over rather short time intervals unless special precautions are taken (see Section II.E.1), and the measurements of current–time curves, called chronoamperometry (CA), are often complicated by additional modes of transport. Therefore, the use of properly shielded electrodes [148] should be considered in chronoamperometric experiments exceeding approximately 1 s. The mathematical formalism for CA has been developed also for the application of ultramicroelectrodes [226] including the transition from nonstationary (increasing diffusion layer thickness) to steady-state transport [198a,227]. The application of CA for monitoring the progress of chemical reactions following the heterogeneous electron transfer is limited by the nature of the process. As it appears from Equation 2.64, the only parameter that may be affected by a follow-up reaction is n. Thus, important mechanisms such as the eC (Scheme 2.2) and  eC(dim) (Scheme 2.3) cannot be characterized by this type of

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Organic Electrochemistry

experiment. However, it has been demonstrated that if the potential is shifted to values insufficient for diffusion-controlled electron transfer, rate constants, for instance, for the eC mechanism, may be determined [228]. Working curves have been calculated relating the apparent number of electrons exchanged napp/n to the dimensionless parameter log(kt), where k is the first-order or pseudofirst-order rate constant for the chemical step: napp it 1/2 = 1/2actual n itdiff.control

(2.65)

An example of the application of the technique is a study of the kinetics of the benzidine rearrangement, which follows the reduction of azobenzene in acidic aqueous ethanol [229]. The approach has been further developed and extended to encompass the automatic recording of three-dimensional i–E–t curves [230]. However, experiments of this kind may be hard to perform with high precision because of difficulties in the accurate control of the potential in the region close to Eo. If, during the chemical step, a product is formed that undergoes further electron transfer at the applied potential, as, for example, in the eCe- and eCeh-type mechanisms (Scheme 2.2), or if the electroactive substrate is regenerated in a catalytic process (Scheme 2.7), the value of napp/n depends on k even under conditions in which the heterogeneous electron transfer is diffusion controlled [66,139]. It is easily understood that napp/n approaches unity when k decreases toward zero, and a higher value, depending on the mechanism, when k increases. The competition between heterogeneous electron transfer (eCe) and electron transfer in solution (eCeh) in the second e step (Scheme 2.2) and the possibility of distinguishing between these two pathways have been analyzed in detail [85,231]. It was concluded that the eCeh pathway dominates over the eCe pathway unless the measurement time is kept below approximately 10 –7 s. The application of CA to determine the rate constants in more complicated reaction schemes, such as, for example, the eCeCe-type mechanism, has been addressed as well [232]. An important extension of CA, double potential step chronoamperometry (DPSCA), involves a second potential step to a potential at which B is converted back to A at the diffusion-controlled rate [1,152,233,234]. The potential–time program and the current–time curve are illustrated in Figure 2.27. The current is measured twice, the first time at t = tf(if ) and the second time usually at t = 2tf(ib). For a simple electron transfer reaction, the ratio −ib/if is independent of tf and has the value 0.2928. 1.5

E (V)

0.5

if 0.0

ib

–0.5

i (arbitrary units)

2tf

tf

1.0

–1.0 –1.5

0

5

10

15

20

25

t (arbitrary units)

FIgURE 2.27 Double potential-step chronoamperometry; potential–time program for a potential step from 0 to −1 V (dotted line and left-hand scale) and current–time curve (full line and right-hand scale).

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143

Techniques for Studies of Electrochemical Reactions in Solution –2

–1

0 log λ

R 1.0

0.5

0.0 log(θ[PhOH]) –5

–4

FIgURE 2.28 Double potential-step chronoamperometry working curve (top scale) for the eCeh mechanism (full line) and experimental data (bottom scale) for the protonation of the anthracene radical anion by phenol in DMF (0.1 M Bu4NBF4). C A* = 1 mM and C *PhOH (mM) = 9.95 (+), 18.2 (•), 40.0 (o), 78.8 (x), 100 (◻), and 200 (*). * . (Reprinted from J. Electroanal. Chem., 147, R at the figure corresponds to RI in this chapter and λ = ktf CPhOH Amatore, C., Gareil, M., and Savéant, J.M., 1–38, Copyright (1983), with permission from Elsevier.)

However, if a chemical reaction is associated with the electron transfer reaction, the ratio −ib/if depends on tf. Working curves based on analytical solutions of the diffusion–kinetics equations have been presented for simple mechanisms [235]. More generally, the working curves may be calculated by digital simulation and are usually presented as the normalized current ratio RI = −ib/(if·0.2928) as a function of log(ktf ) (first-order rate law) or log(kCA* tf ) (second-order rate law; Figure 2.28). The value of k (or kobs) may be determined by fitting the working curve to the experimental data as demonstrated in Figure 2.28 for the protonation of anthracene radical anion by phenol. Alternatively, the rate constant may also be determined directly from the measured value of t1/2, that is, the value of t required to obtain RI = 0.5, in a fashion analogous to that described for DCV earlier (see Equations 2.36 and 2.37) as follows: eC mechanism: kobs =

RR dimerization: kobs =

0.406 −1 s t1/2

0.830 −1 −1 M s t1/2CA*

(2.66) (2.67)

The approximate kinetic window of DPSCA at electrodes with a diameter in the 0.1–2 mm range is given by Equation 2.68, where k is a first-order or pseudo-first-order rate constant. With the application of ultramicroelectrodes, the upper limit can be extended considerably [225]: 0.3 s−1 < k < 300 s−1

(2.68)

DPSCA offers two attractive features, at least. First, since the method includes potential steps to E values at which the heterogeneous electron transfer reaction is diffusion controlled, experiments

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Organic Electrochemistry

may be performed with the iR-compensation circuit of the potentiostat switched off. Second, the perturbation of the reaction layer is very small, which means that even a small excess of reagent relative to substrate ensures pseudo-first-order conditions [97]. A  disadvantage is that DPSCA is a blind technique, and current caused by, for example, long-lived intermediates may erroneously be attributed to regeneration of substrate. In such cases, it may be advantageous to use triple potentialstep CA [237] in which the current flowing after the third potential step reflects whether starting material has indeed been formed after the second step. In general, potential-step techniques should be used only together with a potential-sweep technique, such as CV. Examples illustrating the application of DPSCA include the cleavage of the carbon–halogen bond in radical anions derived from aromatic compounds [238]; the protonation of radical anions derived from aromatic hydrocarbons [97,236,239]; the dimerization of radical anions [119,120,240], radical cations [241], and neutral radicals [101]; and the conversion of the B form to the A form for 10,10′-dimethyl-9,9′-biacridylidene [242].

B.

CHRONOCOULOMETRY AND DOUBLE POTENTIAL-STEP CHRONOCOULOMETRY

A method analogous to CA is chronocoulometry, by which the charge Q instead of the current is monitored as a function of time [1,152,243]. The method has found less widespread use than CA in organic electrochemistry but offers certain advantages when, for instance, an electroactive reactant is adsorbed at the electrode surface [243–247]. Integration of Equation 2.64 with respect to time results in Equation 2.69, which illustrates that a plot of Q versus t1/2 results in a straight line with the intercept Qc, that is, the charge flowing into the interfacial capacitance as a result of the potential step. This quantity contains information about the degree of substrate adsorption at the electrode surface. Q=

2nFACA* ( DAt )1/ 2 + Qc π1/ 2

(2.69)

The application of chronocoulometry for mechanism analysis and evaluation of kinetic parameters is similar to the application of CA and is often based on visual comparison of experimental data with working curves. Double potential-step chronocoulometry [1,152,243] may be used similarly to DPSCA. The working curves now include the charge ratio −Q b/Qf, which takes the value 0.414 for a simple electron transfer reaction. The reductive cyclization of ethyl cinnamate (see Chapter 17) illustrates the use of the technique [248,249].

C. CHRONOPOTENTIOMETRY AND CURRENT-REVERSAL CHRONOPOTENTIOMETRY Chronopotentiometry in its most simple form involves the measurement of the potential of the working electrode as a function of time after the onset of a constant current, i [1,152]. Because of the decrease in concentration of A close to the electrode, the potential changes with time, and when the concentration is sufficiently small, a rapid potential shift is observed to the value of another redox system or to the value for discharge of the solvent-supporting electrolyte solution (Figure 2.29). The time at which this rapid potential shift is observed, called the transition time τ, can, for reversible electron transfer, be calculated from the Sand equation: τ1/ 2 =

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π1/ 2 nFADA1/ 2CA* 2i

(2.70)

145

Techniques for Studies of Electrochemical Reactions in Solution

E

E1/4

τ/4

τ

Time

FIgURE 2.29 Chronopotentiometry; potential–time relationship after application of a constant current. E1/4 is the potential acquired by the electrode at time τ/4.

The complete chronopotentiometric wave is described by Equation 2.71. If t = τ/4, the third term cancels, and Equation 2.71 reduces to Equation 2.72, which defines the quarter-wave potential E1/4 and which is equal to Eo for DA = D B: E = Eo +

RT D RT τ1/ 2 − t 1/ 2 ln B + ln 2nF DA nF t 1/ 2

E1/4 = E o +

RT D ln B 2nF DA

(2.71)

(2.72)

Thus, the quarter-wave potential plays the same role in chronopotentiometry [250] as the half-wave potential E1/2 in polarography (see Section V). The mathematical formalism for chronopotentiometry has been developed also for the application of ultramicroelectrodes [251]. Chronopotentiometry has found only little use in mechanistic organic electrochemistry. This is primarily due to experimental difficulties in the accurate evaluation of the transition time. A solution to this problem includes the application of a convolution procedure [252]. Another extension includes the application of exponential and other current–time functions, and theoretical data for this method are now available for a number of mechanisms [253,254]. When applied to the analysis of kinetic data, chronopotentiometry is most often used in the current-reversal mode, in which the direction of the current is changed after some time tf (Figure 2.30). When only the direction but not the magnitude of the current is changed, the reverse transition time τr is given as follows [255]: τr =

tf 3

tf ≤ τ

(2.73)

Diagnostic criteria for a number of mechanisms have been summarized [256]. The analysis of experimental data is most conveniently carried out with the aid of a working curve (Figure 2.31) relating the parameters tf, τr, and the rate constant k (or kobs). An example of the application of chronopotentiometry is the oxidation of tropylidene [258,259] and thioxanthene [259].

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0

E

E1/4

τ/4

FIgURE 2.30

tf

tf +τr

Time

Current-reversal chronopotentiometry; current–time program and potential–time curve. 0.33 0.30

Tr/tf

0.20

0.10

0.00 0

1

2

3

4

5

kta

FIgURE 2.31 Current-reversal chronopotentiometry working curve for the catalytic eC mechanism. The parameter ta on the figure corresponds to tf in the text. (Reprinted with permission from Testa, A.C. and Reinmuth, W.H., Anal. Chem., 32, 1512–1514. Copyright (1960) American Chemical Society.)

V. POLAROgRAPHy Polarography is the term used for voltammetry with the dropping mercury electrode (DME). The technique has been discussed extensively in several textbooks and reviews [1,152,260–265] to which the reader is referred for details concerning both theoretical problems and practical applications. The electrode (Figure 2.32) was developed early in the twentieth century by Heyrovsky and was the dominating tool in electroanalytical chemistry for several decades. Because of the low oxidation potential of mercury (0.3–0.4 V versus SCE), the DME has been used almost exclusively for the study of reduction processes. Compared with mercury film electrodes, the DME offers the advantage that the electrode surface is continuously renewed. This property reduces undesirable surface effects caused by adsorption.

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h

(+)

FIgURE 2.32

Schematic of a dropping mercury electrode.

A typical current–voltage curve, a so-called polarogram, obtained at the DME as a result of a slow (1–10 mV s−1) linear potential sweep in the negative direction is shown in Figure 2.33. The shape of the polarogram illustrates clearly the expansion of the electrode surface during the lifetime of the individual mercury drops. Two parameters are generally used to characterize the polarogram, the half-wave potential E1/2 and the limiting current ilim. For a simple electron transfer reaction, the height of the current plateau, measured as the mean current during the drop lifetime id, is given by the Ilkovic equation: id = 607nm 2/3t 1/6 DA1/2CA*

(2.74)

i

i E1/2 ilim 1i 2 lim (+) 0 (–)

FIgURE 2.33

0 (V vs. Ref )

E

Polarogram obtained at the dropping mercury electrode.

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(+) 0 (–)

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Organic Electrochemistry

The subscript d indicates that the surface concentration of substrate is dependent only on the diffusion of material to the electrode, m is the mass (mg) of mercury flowing out of the capillary per second, and t is the drop lifetime (s). Polarographic current data are frequently normalized with respect to CA* , m2/3, and t1/6 to eliminate the parameters associated with a particular experiment. The current normalized as in the following equation is called the diffusion current constant I: I=

id = 607nDA1/2 C m 2/3t 1/6 * A

(2.75)

If the electron transfer, Equation 2.1, is reversible, the complete polarographic wave may be described by a modified Nernst equation [260,261]: E = Eo +

RT f DB1/2 RT id − i ln A 1/2 + ln nF fB DA nF i

(2.76)

If i = 0.5 id, the third term in Equation 2.76 cancels, and the expression reduces to the following equation that defines the half-wave potential E1/2: E = E1/2 = E o +

RT f DB1/2 ln A 1/2 nF fB DA

(2.77)

When fA = f B and DA = D B, E1/2 equals Eo. It is seen also that E1/2 is independent of the substrate concentration. Introduction of E1/2 into Equation 2.76 results in the more convenient form, where the constant 0.0591 refers to T = 298 K, as follows: E = E1/2 +

0.0591 RT id − i i −i ln log d = E1/2 + nF i n i

(2.78)

The reversibility of the electron transfer reaction may be tested via this equation, which predicts that a plot of E versus log(( id − i ) / i ) results in a straight line with the slope 0.0591/n V (at T = 298 K) for a reversible redox system. Slopes smaller than 0.0591/n V are observed when the electrode process is quasireversible or irreversible. In the latter case, E and i are related through the following [261,264]: E=

1.349k ot 1/ 2 RT RT i −i for i ≥ id /10 ln + 0.9163 ln d αnF DA1/ 2 αnF i

(2.79)

It should be noted that when Equation 2.79 is used, it is the current at the end of the drop life rather than the mean current that should be measured. In analogy to the reversible case, the half-wave potential can be introduced, Equation 2.80, resulting, finally, in Equation 2.81: E = E1/ 2 =

E = E1/ 2 + 0.9163

1.349k ot 1/ 2 RT ln αnF DA1/ 2

RT i −i 0.0542 i −i = E1/ 2 + ln d log d αnF αn i i

(2.80)

(2.81)

The transfer coefficient α may be obtained from a plot of E versus log((id − i)/i), which should be linear with a slope of 0.0542/(αn) V.

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Both m and t in the Ilkovic equation (Equation 2.74) depend on the height of the mercury reservoir h, and since m is equal to c′h and t is equal to c″/h, where c′ and c″ are constants, we have the relationship between id and h given by Equation 2.82. Thus, the magnitude of a purely diffusion-controlled polarographic wave is proportional to the square root of the height of the mercury reservoir: 1/6

 c′′  id = 607n(c′h)2/3   h

DA1/2CA* = (const) ⋅ h1/2

(2.82)

Polarography has been used to characterize the electrochemical reduction of numerous substrates under aqueous conditions [260–265]. Both E1/2 and ilim may be affected by the proton activity. The polarograms resulting from reduction of anthracene in solvent systems of increasing proton activity illustrate a general trend observed during this type of experiment (Figure 2.34). Under essentially nonaqueous conditions, two reversible one-electron polarographic waves are observed corresponding to the successive formation of the radical anion and the dianion (Figure 2.34, curve a). When the water content increases, the first wave grows at the expense of the second wave, the sum of the two remaining almost constant, until finally at high proton activities, the two waves have merged into one two-electron wave (Figure 2.34, curve e). This behavior can be rationalized in terms of the eCeh mechanism in Scheme 2.2 with X representing a water molecule. The increase in the height of the first wave is related to the increasing importance of the first proton transfer step (ii). In comparison with CV, curve a in Figure 2.34 would correspond to Figure 2.4, whereas curve e in Figure 2.34 would correspond to Figure 2.13 (dotted curve). In purely aqueous solution, the half-wave potential E1/2 for the general reversible reaction shown in Scheme 2.5 is given (at T = 298 K) as follows: pH = 0 E1/2 = E1/2 − 0.0591

m pH n

(2.83)

3

i (μA)

2

1 e

d

c

b

a 0.5 V

Potential

FIgURE 2.34 Effect of increasing proton activity (a through e) on the polarogram of anthracene. For the sake of clarity, the five polarograms have been shifted horizontally. Note also that the oscillations caused by the growth and fall of the mercury drops are not shown. (Hoytink, G.J., in: Delahay, P., ed.: Advances in Electrochemistry and Electrochemical Engineering. Vol. 7. pp. 221–281. 1970. Copyright Wiley-VCH Verlag GmbH & Co. KGaA. Reproduced with permission.)

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Thus, for two reaction schemes commonly met in organic electrochemistry, m = n = 1 and m = n = 2, a plot of E1/2 versus pH should yield a straight line with the slope 59 mV/pH. Numerous examples of the variation of E1/2 with changing pH have been reported [266–268]. Although polarography to some extent has lost importance relative to new powerful techniques, the method deserves not to be forgotten [265]. Polarography is still in many cases a competitive method for the examination of, for example, electron transfer reactions preceded or followed by slow chemical steps or electrochemical reactions associated with adsorption phenomena.

VI. A.

METHODS bASED ON FORCED CONVECTION GENERAL CONSIDERATIONS

It is characteristic for the methods described so far that the measurements are carried out in a quiet solution and that the transport of material to and from the working electrode is assumed to be governed by diffusion only (except in solution with very low concentration of supporting electrolytes). It has been mentioned also that undisturbed diffusion cannot be maintained over longer periods of time unless the working electrode is properly shielded. Another way to suppress the effects caused by natural convection is to use a rotating disk electrode (RDE). Owing to the rotation of the electrode, a strong convective force is imposed on the solution (see Section VI.B.), and in comparison, other forces that may influence the solution transport are small and can be neglected. Furthermore, the diffusion layer thickness for a rotating electrode is stagnant and small compared to that of a stationary electrode, and taken together, these characteristics result in an experimental setup that is relatively insensitive to the nature of the surroundings. However, this advantage has to be paid for by the drawback that the RDE, when used in organic electrochemistry, often suffers from electrode fouling presumably caused by the deposition of insulating films, a problem related to the high material conversion at the electrode surface. Thorough discussions of the applications of RDEs and the associated mathematics are found in several texts [1,269–272].

B.

THE ROTATING DISK ELECTRODE

The electrode tip of an RDE system may consist of a cylindrical piece of conducting material (usually Pt, Au, or glassy carbon), which together with a wire or a steel rod for electrical contact is sealed into an insulating shaft (Figure 2.35). When the electrode is rotated, an adjacent thin layer of the solution dragged by the electrode surface acquires the tangential velocity of the electrode. Because of the centrifugal force, this is superimposed by a radial velocity, resulting in an overall flow pattern at the electrode surface, as illustrated in Figure 2.36. The liquid thrown out horizontally is replaced through an upward vertical flow (Figure 2.35). The mathematical treatment on which the equations given in this section are based is valid only as long as the fluid flow is laminar. At very high rotation rates, the flow becomes turbulent, and at very low rotation rates, natural convection begins to play an unwanted role. The two extremes can be given in terms of the Reynolds number Re, which for the RDE is defined in Equation 2.84, where r is the radius (cm) of the RDE, including the insulating shaft (!), ω is the angular rotation rate (rad s−1), and ν k is the kinematic viscosity coefficient (cm2 s−1). Laminar flow can be maintained for Re values between 10 and 104, both values being approximate. For a typical commercial RDE, these extremes limit the useful range of rotation rate to values roughly between 0.2 and 200 rev s−1. Re = r 2 ων −k1

(2.84)

The tangential and radial velocities decrease with increasing distance from the electrode surface, and at some critical distance δo, it is sufficiently accurate to consider the laminar upward motion as the only mode of transport. The liquid layer in which it is necessary to take into account both the

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Electrical contact Insulating shaft

Conducting material

FIgURE 2.35

Schematic of a rotating disk electrode. The streamlines indicate the flow below the disk.

FIgURE 2.36

Streamlines at the electrode surface of a rotating disk electrode.

tangential and radial velocities is called the hydrodynamic boundary layer, and its thickness is given approximately by Equation 2.85. Thus, the value of δo varies between 2.7 and 0.08 mm for rotation rates in the range 0.2 and 200 rev s−1 assuming the value 0.01 cm2 s−1 for νk. δo = 3ν1k/ 2ω−1/ 2

(2.85)

When an electrochemical process takes place at the electrode surface, a concentration gradient develops near the surface, resulting in diffusion as an additional mode of mass transport. The liquid layer in which the transport by diffusion is comparable to the convectional motion is called the diffusion boundary layer, and its approximate thickness δ is given by Equation 2.86, corresponding to approximately 5% of δo with viscosity and diffusion coefficient in typical ranges. δ = 1.61DA1/3ν1k/6ω−1/ 2

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(2.86)

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Organic Electrochemistry

i

A

B

C 2

ilim

1 0

E0

E

FIgURE 2.37 Current–voltage curves for a rotating disk electrode: (1) without and (2) with substrate added. The regions indicated are as follows: (A) control by electron transfer; (B) mixed control by electron and mass transfer; (C) control by mass transfer.

Both δo and δ have the same value over the entire electrode surface, which has given rise to the description of the electrode as a uniformly accessible surface. The concepts of a hydrodynamic and a diffusion boundary layer have no theoretical significance as such but serve mainly to provide a suitable model for the hydrodynamic conditions related to the rotating electrode. At constant rotation speed, the current–voltage curve (Figure 2.37) for a simple reversible electron transfer reaction, Equation 2.1, is similar to a polarographic wave, and the half-wave potential E1/2 for the current–voltage curve obtained by an RDE is defined in the same way as in polarography (Section V). The curve can be considered as consisting of three regions, a, b, and c, depending on the value of the applied potential compared to Eo. For the general case where the process is under the mixed control of mass and electron transfer (region b in Figure 2.37), the current is given as follows [273,274]: i=

nFACA* DA nFACA* DA = −1/ 2 1.61D ν ω + ( DA /ks,f ) δ + ( DA /ks,f ) 1/ 3 1/ 6 A k

(2.87)

For a typical value of δ ≈ 10 −3 cm, the term DA /ks,f can be neglected if ks,f > 10 −1 cm s−1. The current is now controlled solely by the mass transfer to the electrode (region c in Figure 2.37), and the equation reduces to the so-called Levich equation, Equation 2.88, which may be used together with the Cottrell equation, Equation 2.64, in the experimental determination of δ [275]: ilim = 0.62nFACA* DA2 /3ν −k1/6ω1/ 2 =

nFAC A* DA δ

(2.88)

An important consequence of this equation is the prediction that ilim varies linearly with ω1/2, a crucial test of the quality of the RDE. On the other hand, if ks,f < 10 −4 cm s−1, the term DA /ks,f is large compared to δ, and for this situation, in which the current is controlled by the rate of the electrochemical process (region a in Figure 2.37), Equation 2.89 is valid: i = nFACA* ks,f

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(2.89)

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153

These relations have been discussed and experimentally tested [273]. Values of ko up to approximately 0.1 cm s−1 may be measured using the RDE. By inspection of Equations 2.87 through 2.89, it is seen that measurement of the rotating disk current may provide information about either n, CA* , DA, νk, or ks,f when the remaining parameters are known. Accuracy in determination of n values depends mainly on how accurately the value of DA for the compound in question is known. For the majority of organic compounds, diffusion coefficients in the common organic solvents have not been measured, but comparison with a suitable compound of similar structure for which both n and D are known can often solve the problem (assuming the same value of D for the two substrates; see also Section VIII for methods to determine D). Since DA appears only to the power of 2/3 in Equation 2.88, the resulting n value is only moderately dependent on the uncertainty of the DA estimate. The measurement of diffusion coefficients by means of the RDE technique has been the subject of a number of papers [271,276] and need not be treated further at this place. A major field of application of the RDE in organic electrochemistry is elucidation of the mechanistic details for reactions following the heterogeneous electron transfer reaction. One way of using the RDE in kinetic work is based on the relationship between ilim and ω. In the absence of kinetic complications, a plot of ilim versus ω1/2 should be a straight line (Equation 2.88). However, if the electron transfer reaction is followed by a chemical step, the linear dependence of ilim on ω1/2 may change to a curved relationship. For the eCe mechanism (Scheme 2.2), the relationship between the 1 are apparent number of electrons, napp, ilim, and ω is given by Equation 2.90 [277], where n and ilim the number of electrons and the current observed for the first electron transfer step only, that is, in the absence of the chemical reaction. Other mechanisms have been treated as well [271,278,279]. napp ilim ( −0.834 ν1k/3k /( D1A/3ω)) = 1 = 2−e n ilim

(2.90)

It is easily seen that for relatively fast chemical steps where k ≫ ω, the exponential term is close to zero, and the current corresponds to the exchange of two electrons. On the other hand, if k = 0, the current will correspond to a one-electron process. Consequently, for intermediate values of k, the apparent number of electrons varies between 1 and 2. Equation 2.90 predicts that napp tends toward smaller values with increasing ω. The working curve showing the relationship between napp and ω (ω1/2) can be calculated from Equation 2.90 for any value of k provided that the remaining parameters are known. Working curves at different k values for the eCeh mechanism (Scheme 2.2) are shown in Figure 2.38. The range of measurable first-order rate constants is given as follows: 0.1 s−1 < k < 103 s−1

(2.91)

Typical studies, in which the RDE has been used, include the dimerization of the triphenylamine radical cation [280] and the cyclization of the tetraphenylethylene dication [281]. Another approach includes measurements of E1/2 relative to that for the simple electron transfer reaction [282–284]. For instance, for the eCeh mechanism, ΔE1/2 is given by Equation 2.92, and it is seen that a plot of ΔE1/2 versus log ω is predicted to be linear with the slope −2.303RT/(2nF) (= −29.5 mV at n = 1 and T = 298 K). The value of k can easily be determined from this plot.  D  2.303RT log  A  nF  ν k 

1/ 6

∆E1/ 2 = −

ω   + log 2 0.643 k 

(2.92)

An example is shown in Figure 2.39 for the reduction of fluorescein in aqueous solution at pH 9.5–9.7 [282].

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1.9 1000 s–1 1.8 300 s–1 1.7

100 s–1

nobs

1.6

1.5

1.4

30 s–1

1.3

1.2

10 s–1

1.1

3 s–1

1.0 5

10

15

20

ω1/2

FIgURE 2.38 Rotating disk electrode working curves for the eCeh mechanism at different values of the rate constant k for the chemical step.

The applications of the RDE presented so far all relate to operation under steady-state conditions. However, it has been shown that the possibility to discriminate between closely related mechanisms, such as the eCe and the eCeh, may be improved considerably by using a potential-step technique together with the RDE. The reader interested in such details is referred to the original literature [279,285,286].

C.

THE ROTATING RING-DISK ELECTRODE

At the RDE, the electrode products are thrown horizontally away from the electrode surface and thus escape further detection. A solution to the need for additional information in RDE experiments is the rotating ring-disk electrode (RRDE) [287]. The electrode consists of a conducting disk and a conducting ring in a concentric arrangement separated by insulating material (Figure 2.40). By proper adjustment of the ring potential, initial electrode products or products from followup reactions may be monitored as a function of the angular rotation rate. The parameter most frequently used in RRDE experiments is the collection efficiency N, which relates the ring current ir to the disk current id as follows: N=

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ir id

(2.93)

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Techniques for Studies of Electrochemical Reactions in Solution

E1/2 (V) vs. SCE

–1.150

–1.140

–1.120

–1.110 –0.4

–0.2

0

0.2

0.4

log10 (ω/Hz)

FIgURE 2.39 Variation of the half-wave potential for the reduction of fluorescein with the rotation speed for the current–voltage curves obtained at pH 9.68 (⬩) and 9.51 (x). (Compton, R.G., Harland, R.G., Unwin, P.R., and Waller, A.M., J. Chem. Soc., Faraday Trans. 1, 83, 1261–1268, 1987. Reproduced by permission of The Royal Society of Chemistry.)

r1

r2

r3

FIgURE 2.40 Schematic of a rotating ring-disk electrode.

These observables usually have opposite signs and should be measured at potentials where mass transport is rate determining. The value of N is less than unity, even for a simple electron transfer reaction, because only a fraction of the molecules generated at the disk electrode reaches the ring owing to diffusion away from the electrode. The theoretical value of N depends on the geometry of the electrode system. Of special importance is the size of the gap between the two electrodes r 2 − r1 (see Figure 2.40). The value of N typically varies between 0.45 and 0.2 [288] and is independent of ω for a simple electron transfer reaction. If the initial electrode product B undergoes a chemical reaction, the observed ring current, and therefore N, drops to a lower value. The fraction of the B molecules that escape detection at the ring depends on two factors, the rate of the chemical step and the time interval between the generation

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and detection of B, that is, on ω. Consequently, it can be expected that for a given chemical reaction, the value of N increases with increasing ω and in the limit approaches the value for the simple electron transfer reaction. Working curves for the determination of rate constants can be constructed by analogy to those already described for the RDE, this time relating N to ω or an expression containing this parameter. Because of the complexity of the mathematics associated with the calculation, digital simulation has been extensively used for this purpose [288–290]. Comparison of analytical and numerical solutions for a number of cases demonstrated satisfactory agreement between the two approaches [291]. If the mechanism is known, the rate constant can be obtained from the appropriate working curve, which for the RRDE usually relates the collection efficiency to one of the two parameters given in the following [288]: XKT = (0.51)−2 /3 kω−1ν1k/3 DA1/3 (first  order rate law)

(2.94)

XKTC = (0.51)−2 /3 kCA* ω−1ν1k/3 DA1/3 (second  order rate law)

(2.95)

However, for a specific case, it may be more convenient to use a set of working curves, similar to those shown for the RDE, that relates N to ω1/2 (Figure 2.41). Application of the RRDE as a diagnostic tool to explore the mechanism of an unknown reaction pathway relies, as most methods do, on the differences between the different working curves. For the RDE, the working curves for different mechanisms in many cases do not deviate sufficiently to allow a distinction between two possible pathways [282]. The presence of the ring electrode provides extra information, and thus the RRDE is a more sensitive instrument for mechanism analysis. An example is the distinction between the RR- and the RS-dimerization mechanisms for CS2 radical anion [292]. It has been pointed out, however, that even with this electrode setup, the working curves do not allow the unequivocal discrimination between, for example, eCe-type (Scheme 2.2) and catalytic mechanisms (Scheme 2.7) [288] or the different dimerization mechanisms (Scheme 2.3) [289]. A solution to this problem has been found, which 0.2 s–1

0.5 s–1

1 s–1

2 s–1 5 s–1

0.15 10 s–1

20 s–1

0.00

50 s–1

N

0.10

0 2.5

5.0

7.5

10.0

12.5

15.0 17.5

20.0

22.5 25.0

27.5

30.0

ω1/2

FIgURE 2.41 Rotating ring-disk electrode working curves for the eCe mechanism at different values of the first-order rate constant k for the chemical step. The value of N corresponding to k = 0 is 0.184 for this particular electrode.

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takes advantage of the possibility of varying the relative fluxes of A and B from the disk via the current density [289,293–295]. The RRDE technique can provide quantitative information for reactions having rate constants in the ranges given by Equations 2.96 and 2.97. The limits are dependent on the dimensions of the RRDE, in particular the size of the gap between the ring and the disk.

VII.

10 −1 s−1 < k < 103 s−1 (first-order rate law)

(2.96)

10 M −1 s−1 < k < 108 M −1 s−1 (second-order rate law)

(2.97)

METHODS FOR DETERMINATION OF THE NUMbER OF ELECTRONS TRANSFERRED PER MOLECULE OF SUbSTRATE

An important parameter in all kinds of electrochemical work, including mechanism analysis, is the number of electrons n transferred per molecule during the electrochemical reaction. Values of n may be obtained by almost any electrochemical technique provided that a number of other parameters are known as illustrated, for example, by the relations between n and i given in Equations 2.13, 2.61, 2.64, 2.74, and 2.88. However, it is the exception rather than the rule that all the necessary parameters are known for a particular experiment. This applies especially to the effective electrode area A and diffusion coefficient DA, which is known only for few substances in the solvent systems commonly used in electrochemical studies (see also Section VIII). A simple solution to this kind of problem is based on the quantitative comparison of, for instance, LSV and CA [296]. From Equations 2.13 and 2.64, it is easily seen that the ratio R = (ip/ν1/2)/(it1/2) is given by the following simple expression, since all other parameters cancel: R=

(ip / ν1/ 2 ) = 4.92n1/ 2 it 1/ 2

(2.98)

Thus, it is possible to estimate n values without the prior knowledge of DA and A, provided the same working electrode is used for both experiments. The experimental values of R for systems with known n were found to agree with the theoretical values (4.92 for n = 1 and 6.96 for n = 2) within 6%. This approach for the determination of n has later been discussed in more detail elsewhere [297]. A different way of determining n values is based on the measurement of the amount of charge necessary for the exhaustive electrolysis of a known amount of substrate. This type of experiment, traditionally called coulometry, may be carried out either at constant potential or at constant current. During coulometry at constant potential, the amount of charge is obtained by integration of the current–time curve (Figure 2.42). The advantage of potential control during the coulometry experiment has to be paid for by a rather long electrolysis time. The end point, i = 0, is in principle not reached until infinite time has elapsed. In practice, however, the electrolysis is stopped when the current has decayed to a few percent of the initial value. The error introduced in this way can be neglected for all practical purposes. Coulometry at constant current is often considered as being less attractive than coulometry at constant potential. Control of the current rather than the potential has, however, a number of advantages. First, the charge consumed during the reaction is directly proportional to the electrolysis time, and second, simpler electronic equipment may be used. It has been demonstrated [299] that when the current density is low, the potential of the working electrode stays almost constant until 90% of the substrate is consumed. The total conversion of, for example, 0.1 mmol of substrate at 25 mA requires an electrolysis time of 6.44 min for n = 1. The change in concentration during electrolysis is conveniently followed by LSV or CV. Cyclic voltammograms of a solution containing

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Time

FIgURE 2.42 The current i as a function of the time t in constant potential electrolysis. The current is given by i = it=0·10 −kt, where k is a constant depending on the construction of the electrochemical cell. (From Lingane, J.J., J. Am. Chem. Soc., 67, 1916, 1945.)

0.1 mmol 3,4-diphenyl-1,2-dithiolylium ion [300], after constant current (25 mA) reduction in 0, 1, 2, 3, 4, 5, and 6 min, respectively, are shown in Figure 2.43. Extrapolation of the current to ip = 0 gives time for complete conversion, which in the present case was very close to the time, 6.44 min, predicted for n = 1. Under conditions where the primary electrode product undergoes a slow chemical reaction, that is, t1/2 is of the order of seconds, the value of n determined by a relatively fast technique like LSV may differ from that obtained by a slow experiment like coulometry. This type of behavior was observed in the anodic oxidation of 2,3,5,6-tetraphenyl-1,4-dithiin in MeCN [301]. During CV, the reversible oxidation to the radical cation is observed. However, when constant-current

I

I

(+)

(+) 0

0 (–)

(–) 0

1

2

3

4

5

6 Ph

S

S +

2

S

S

2e–

H Ph

Ph

Ph

H S

Ph

S

Ph

FIgURE 2.43 Constant-current coulometry (i = 25 mA) of 3,4-diphenyl-1,2-dithiolylium ion (0.1 mmol). Cyclic voltammograms, horizontally displaced, are shown after reduction for 0, 1, 2, 3, 4, 5, and 6 min, respectively.

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(+) 0

(+) 0

(–)

(–)

0

1

2

3

4

5

6

Electrolysis time (min) Ph

S

Ph

Ph

Ph S

FIgURE 2.44 Constant-current coulometry (i = 50 mA) of 2,3,5,6-tetraphenyl-1,4-dithiin (0.1 mmol). Cyclic voltammograms, horizontally displaced, are shown after oxidation for 0, 1, 2, 3, 4, 5, and 6 min, respectively.

coulometry was carried out as described earlier, this time at i = 50 mA, 6.44 min was required to oxidize completely 0.1 mmol of the substrate to a product electroinactive in the potential region of interest, indicating an overall two-electron process (Figure 2.44). Thus, apparently contradictory results may be obtained due to the difference in time scale between the two types of experiment. The equipment necessary for the determination of n by coulometry as described earlier does in principle not differ from that used for a microscale preparative electrolysis. Therefore, in addition to the determination of n, it is possible also to investigate the voltammetric properties of the product and, for example, whether the starting material may be regenerated by electrolysis after the direction of the current has been changed.

VIII.

METHODS FOR DETERMININg THE DIFFUSION COEFFICIENT OF ELECTROACTIVE MOLECULES

Similar to the determination of n (Section VII), diffusion coefficients may be estimated from various electrochemical methods discussed in this chapter, provided some other parameters, in particular the electroactive area of the electrode A, the bulk concentration of the compound C *, and n, are known. Thus, for example, Equation 2.13 can be used for the determination of D from cyclic voltammetric peak currents. As mentioned already, however, not always all of these data are available. While A is accessible from calibrating experiments with compounds of known n, D, and C * (often ferrocene can be used for this purpose, with n = 1 and D having been reported in a number of commonly used nonaqueous solvents [302]), it is often necessary that n and D have both to be determined for a given system. One alternative is to use a combination of macro- and ultramicroelectrode data. In the former case, Equation 2.13 is valid, while in the latter, Equation 2.61 applies. Since n and D are related

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by different mathematical functions in these equations, both parameters can simultaneously be determined. It has been shown that combining steady-state (ultramicroelectrode) data with CA (see below) may result in minimum uncertainties [297]. In some cases, D and the electrode radii (or electroactive areas) [302,303], or D and C * [304] were estimated. A simultaneous determination of D and A [302] minimizes calibration errors in electrode size. An additional problem with CV data is encountered resulting from the fact that the numerical constant 0.4463 in the Randles–Ševčík equation (2.13) is strictly valid only for reversible electrode processes without any complication by coupled kinetics. Furthermore, if we consider a reaction with two consecutive one-electron steps, Equation 2.16, with overlapping waves, this constant critically depends on E2o − E1o. Only at the extreme cases of fully separated waves (Figure 2.4; two times n = 1) or strong potential inversion [60,305] (Chapter 11; n = 2) Equation 2.13 can be applied. For E2o − E1o = 35 mV, the voltammogram corresponds to two superimposed nonseparable, one-electron curves, with ip = 2·ip (n = 1), and consequently, D is also accessible. Intermediate cases, however, do not allow easy interpretation with respect to D. This is an intrinsic problem of the fact that CV uses a potential sweep. Thus, CA has been a common method for such experiments, because the step is usually performed into the limiting current region of the potential [303,306]. The planar (Cottrell; Equation 2.64) and hemispherical diffusion conditions could be realized in chronoamperograms at short and long times, respectively [307]. Several papers deal with the CA-based estimation of D without the knowledge of the exact concentration [297,303b,306–308]. Convolution voltammetry (see  Section II.F.1) also proved to be advantageous in respect even in room-temperature ionic liquids where diffusion is inherently slow owing to the high electrolyte viscosity [309]. In still other cases, D was determined by non-electrochemical techniques, for example, pulse gradient spin-echo nuclear magnetic resonance experiments [302,310]. Thus, D follows from a nonelectrochemical method, and any errors in the current data and their interpretation can be avoided. The technique was also used in ionic liquids of higher viscosity than common typical organic solvents with corresponding lower D [311]. Estimation by comparison with similar molecules [312] (see however, the cautious remarks in Reference 313) was also used. Diffusion coefficients have also been estimated from molecule sizes [314].

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301. Hammerich, O. Unpublished results. 302. Janisch, J.; Ruff, A.; Speiser, B.; Wolff, C.; Zigelli, J.; Benthin, S.; Feldmann, V.; Mayer, H.A. J. Solid State Electrochem. 2011, 15, 2083–2094. 303. (a) Baur, J.E.; Wightman, R.M. J. Electroanal. Chem. 1991, 305, 73–81; (b) Jung, Y.; Kwak, J. Bull. Korean Chem. Soc. 1994, 15, 209–213. 304. Collinson, M.M.; Zambrano, P.J.; Wang, H.; Taussig, J.S. Langmuir 1999, 15, 662–668. 305. Bard, A.J.; Faulkner, L.R. Electrochemical Methods, 2nd edn.; Wiley: New York, 2001; p. 243ff. 306. (a) Baranski, A.S.; Fawcett, W.R.; Gilbert, C.M. Anal. Chem. 1985, 57, 166–170; (b) Denuault, G.; Mirkin, M.V.; Bard, A.J. J. Electroanal. Chem. 1991, 308, 27–38; (c) Han, L.-M.; Suo, Q.-L.; Luo, M.-h.; Zhu, N.; Ma, Y.-Q. Inorg. Chem. Commun. 2008, 11, 873–875. 307. (a) Winlove, C.P.; Parker, K.H.; Oxenham, R.K.C. J. Electroanal. Chem. 1984, 170, 293–304; (b) Kulesza, P.J.; Faulkner, L.R. J. Am. Chem. Soc. 1993, 115, 11878–11884. 308. (a) Kakihana, M.; Ikeuchi, H.; Satô, G.P.; Tokuda, K. J. Electroanal. Chem. 1980, 108, 381–383; (b) Kakihana, M.; Ikeuchi, H.; Satô, G.P.; Tokuda, K. J. Electroanal. Chem. 1981, 117, 201–211; (c) Lawson, D.R.; Whiteley, L.D.; Martin, C.R.; Szentirmay, M.N.; Song, J.I. J. Electrochem. Soc. 1988, 135, 2247–2253; (d) Whiteley, L.D.; Martin, C.R. J. Phys. Chem. 1989, 93, 4650–4658; (e) Evans, R.G.; Klymenko, O.V.; Saddoughi, S.A.; Hardacre, C.; Compton, R.G. J. Phys. Chem. B 2004, 108, 7878– 7886; (f) Guo, Y.; Kanakubo, M.; Kodama, D.; Nanjo, H. J. Electroanal. Chem. 2010, 639, 109–115. 309. Bentley, C.L.; Bond, A.M.; Hollenkamp, A.F.; Mahon, P.J.; Zhang, J. Anal. Chem. 2014, 86, 2073–2081. 310. (a) Goldsmith, J.I.; Takada, K.; Abruña, H.D. J. Phys. Chem. B 2002, 106, 8504–8513; (b) Sun, H.; Chen, W.; Kaifer, A.E. Organometallics 2006, 25, 1828–1830. 311. Grills, D.C.; Matsubara, Y.; Kuwahara, Y.; Golisz, S.R.; Kurtz, D.A.; Mello, B.A. J. Phys. Chem. Lett. 2014, 5, 2033–2038. 312. (a) Valencia, D.P.; González, F.J. Electrochem. Commun. 2011, 13, 129–132; (b) Valencia, D.P.; González, F.J. J. Electroanal. Chem. 2012, 681, 121–126. 313. Parker, V.D. Electrochim. Acta 1973, 18, 519–524. 314. Ruiz Abad, D.; Henig, J.; Mayer, H.A.; Reißig, T.; Speiser, B. Organometallics, 2014, 33, 4777–4783.

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3

In Situ Spectroelectrochemistry of Organic Compounds Peter Rapta, Evgenia Dmitrieva, Alexey A. Popov, and Lothar Dunsch

CONTENTS I.

Development and the Recent State of In Situ Spectroelectrochemistry ............................... 169 A. Scope of Spectroelectrochemistry.................................................................................169 B. In Situ Spectroscopic Methods Used with Electrochemistry ........................................170 C. Triple In Situ UV–Vis–NIR/ESR Spectroelectrochemistry ......................................... 172 1. Method ...................................................................................................................173 2. LIGA Electrodes in Voltammetric and Spectroelectrochemical Studies ...............174 3. In Situ UV–Vis–NIR/ESR Spectroelectrochemistry of Conducting Polymers......176 4. In Situ UV–Vis–NIR/ESR Spectroelectrochemistry at Variable Temperatures ... 177 II. Spectroelectrochemistry of Organic Compounds in Solution or as Solid Layers on the Electrode Surface.................................................................................................................. 179 A. Empty Fullerenes ..........................................................................................................179 B. Derivatives of Fullerenes ..............................................................................................179 C. Endohedral Fullerenes ..................................................................................................181 D. Dimerization Reactions ................................................................................................ 183 E. Conducting Polymers .................................................................................................... 186 References ...................................................................................................................................... 187

I. A.

DEVELOPMENT AND THE RECENT STATE OF IN SITu SPECTROELECTROCHEMISTRy SCOPE OF SPECTROELECTROCHEMISTRY

Although a large variety of electrochemical methods is nowadays available [1–6] (see also Chapter 2), there is no way to get any structural information about species formed on the electrode upon redox events. Modern electrochemistry is intensively focused on electrode reaction mechanisms, and spectroelectrochemistry helps us to get a detailed picture about the intermediates and products formed as a result of the heterogeneous electron transfer reaction [7–11]. Organic compounds that can be easily and reversibly oxidized and/or reduced received a lot of interest in recent times. The redox processes of such redox-active compounds can lead to different redox states where each of these species exhibits different chemical and spectroscopic properties, which are very meaningful for the application purposes of certain structures. Additional information from various spectroscopic methods, which give direct evidence for the chemical species in the complex redox reactions pathway, frequently helps to find a plausible reaction mechanism for redox processes accompanied by electron transfers [12]. 169

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Spectroelectrochemistry deals with the theory and application of spectroscopic methods in electrochemistry as a bridge between electrochemistry and spectroscopy. In general, spectroscopic methods are applied in electrochemical systems under the condition that the electromagnetic radiation does not change the electrochemical reaction and the equilibrium at the electrode. Electrochemistry is focused on processes at interphases, while spectroscopic methods are often dealing with bulk properties of different materials. Modern research in spectroelectrochemistry is characterized by a combination of different spectroscopic methods for a detailed study of electrode reaction mechanisms or complex electrode systems. In  situ spectroelectrochemistry includes all spectroscopic measurements at a working electrode under electron transfer. Spectroscopic methods can be applied both to the solid and to the liquid phases at the electrode–electrolyte interface. Although a huge variety of methods arose in the last decades concerning spectroscopic methods applied in electrochemistry, there are some spectroscopic methods that are preferred in their use in spectroelectrochemistry.

B. IN SITU SPECTROSCOPIC METHODS USED WITH ELECTROCHEMISTRY The simultaneous use of ultraviolet (UV)–visible, infrared (IR), electron spin resonance (ESR), and electrochemical methods is now widely accepted and is well developed as a powerful tool in the identification of paramagnetic intermediates or other reaction products [13], and there have been published already several monographs in the field. UV–vis–near IR (NIR) spectroscopy is the most applied method in spectroelectrochemistry [14–21]. The electronic structure of molecules and the changes induced by charge transfer reactions are studied by absorption spectroscopy in the range from NIR through visible to the UV region with photon energies from below 1 up to 6.5 eV. This completely covers the field of electronic transitions in molecules and crystal lattices. For instrumental reasons, these three energy ranges are often applied in the same apparatus and are also considered in a combined treatment of the spectroscopic data. These electronic and vibronic transitions are preferably measured in the absorption mode, while the emission of light from excited single molecules and solids is monitored by luminescence spectroscopy [22–30]. An increased understanding of redox processes and various complex reactions in all kinds of organic systems associated with paramagnetic intermediates has been strongly promoted by the application of in  situ ESR spectroelectrochemical techniques [31,32]. ESR spectroscopy is both very sensitive and delivers detailed structural information on paramagnetic species. A variety of ESR spectroelectrochemical cells having different design and construction have been developed concerning both qualitative and quantitative works as well as for measurements at different temperatures as illustrated for several basic cell designs in Figure 3.1 [33–48]. In the electrochemical formation of radicals occurring in one-electron transfer processes of organic species, in situ ESR spectroscopy has been applied successfully in the elucidation of reaction mechanisms. If the electrode reaction product is diamagnetic, in situ UV–vis–NIR spectroscopy is widely recognized as the most valuable spectroscopic method. This method, which can be used both in reflection and in transmission, has been applied to mechanistic studies, although the structural information so obtained is not as detailed as that in ESR spectroscopy. The field of spectroelectrochemistry was broadened by the introduction of optically transparent electrodes (OTEs) for absorption spectroscopy in transmission [49], the introduction of the attenuated total reflection (ATR) technique [50], and the use of thin-layer cells for IR spectroelectrochemistry [51]. OTEs mostly consist of a transparent material (glass, quartz) covered with a thin noble metal film (Pt, Au) or oxides of Sn and In. Alternatively, the minigrid electrodes made from a wire mesh can be used as OTE where the transparency is due to the holes in the metal grid. An optically transparent thin-layer electrode within the thin-layer cells is already widely used due to a small volume of the electrolyte solution needed and due to the fast conversion of the redox system in solution under thin-layer conditions [52–57].

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RE

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FIgURE 3.1 Illustrative overview of some basic ESR spectroelectrochemical cells. (a) Flat cells: (a1 and a2) from Reference 33, (a3) from Reference 34, (a4) from Reference 35, (a5) from Reference 36. (b) Flow-through channel flat cell from Reference 37. (c) Tubular flow-through cells: (c1) from Reference 38 and (c2) from Reference 39. (d) Stationary tubular cells: (d1) from Reference 40, (d2) from Reference 41, (d3) from References 42 and 43, (d4) from Reference 43. (e) Flow-through tubular cells: (e1) from Reference 44, (e2) from Reference 45, (e3) from Reference 46, (e4) from Reference 47. (f) Column cell from Reference 48.

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Vibrational spectroscopy is nowadays widely used as a spectroelectrochemical technique resulting in detailed information on molecular structures [58–61]. Raman spectroscopy provides complementary data regarding non-IR active vibrational modes. A breakthrough in Raman spectroelectrochemistry was the discovery of surface-enhanced Raman patterns at silver electrodes [62]. Together with the availability of the laser Raman technique, this vibrational spectroscopy was since then a standard method in in situ spectroelectrochemistry. In situ Raman spectroelectrochemistry was proven as a powerful tool in the study of redox processes of carbon nanostructures [63]. Nuclear magnetic resonance (NMR) spectroelectrochemistry of molecular structures was made accessible in 1975 by a spinning cell construction to get well-resolved NMR spectra, but it took three more decades to develop a more suitable in situ NMR cell with high sensitivity for a broad range of in situ NMR spectroelectrochemical studies [64–70]. Use of an electrochemical cell within the detection area of an NMR spectrometer enables the structural identification of electrochemically generated species. Static in  situ NMR spectroelectrochemical cells require a low solution volume, and several working electrode designs have been already published including Sb–SnO2 deposited on the inner surface of the NMR tube [65], NMR tubes coated with gold [67,69], or working electrodes based on carbon fibers operating under a large variety of frequencies [68,70]. In general, absorption spectroscopy in the UV–vis–NIR range remains the main spectroscopic method for the characterization of the electronic structure, and the changes occurring from charge transfer reactions are monitored by electrochemistry using in situ methods. Although each spectroscopic technique is considered with respect to its advantages in electrochemical research, it is obvious that both vibrational (IR and Raman) and optical absorption spectroscopies are preferred in electrochemistry due to their valuable chemical information. Fourier transform infrared (FTIR) spectroscopy is the most widespread method and is based on effective cell constructions. On the other hand, NMR spectroelectrochemistry delivers useful structural details of molecules and solids present within the electrochemical system. The advantage of these methods in spectroelectrochemistry is their applicability in the study of an electrochemical system, as both the solid and liquid phases are accessible to a detailed spectroscopic study. However, due to the variety of the investigated systems as well as of the experimental conditions used (e.g., nonaqueous and aqueous solutions, homogeneous and heterogeneous systems and polymers, and nontransparent and optically transparent electrodes), it was necessary to construct special spectroelectrochemical cells for each system individually. In in situ spectroelectrochemical methods, the measurement within an electrochemical experiment at the same electrode system is realized. Advantages of in situ methods in electrochemical studies include mainly a direct access to kinetic data of electrode reactions and qualitative and quantitative information on the state of the interface at electrochemical conditions. Simultaneous acquisition of data from different methods in one single experiment and separation of the faradaic and non-faradaic parts of an electrochemical reaction by a quantitative identification of the reaction products at the electrode also contribute to the power of these methods. However, there are also some disadvantages of in situ methods in electrochemical studies such as the low concentration of the reaction products at the phase boundary and the high time consumption with respect to the preparation of the experimental setup, the simultaneous data acquisition, and the evaluation of all data. Special cells are required, the type and size of which are adjusted to the requirements of the spectroscopic method, which might contradict the requirements of the electrochemical method in terms of electrode geometry, electrolyte composition, and volume. The selection of solvents, supporting electrolytes, and the electrode materials is limited by the requirements of the spectroscopic method.

C.

TRIPLE IN SITU UV–VIS–NIR/ESR SPECTROELECTROCHEMISTRY

In combination with ESR and UV–vis–NIR spectroscopy, cyclic voltammetry (CV) is an important analytic method for estimating the ability of organic compounds to lose or to accept electrons and to elucidate the chemical structure and the electronic nature of the reduced or oxidized species

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generated in course of the corresponding electron transfer processes. The development of an optical ESR cavity opened the route to a simultaneous application of both ESR and UV–vis–NIR spectroscopy in a single in situ spectroelectrochemical technique (triple method) for studies at the same working electrode [71]. In this way, both the paramagnetic and diamagnetic structures can be followed in electrode reactions at the same working electrode. Additional information from optical spectroscopic methods (UV–vis–NIR), which directly refers to both the diamagnetic and paramagnetic chemical species in complex redox reaction pathways, support the search for a plausible reaction mechanism. 1. Method In order to obtain structural information on both the paramagnetic and diamagnetic species in an electrode reaction, it is required to combine the spectroscopic methods mentioned earlier, that is, to measure both in situ UV–vis–NIR and ESR spectra simultaneously during the electrochemical experiment (Figure 3.2). Simultaneous in  situ measurements by both ESR and UV–vis spectroscopy have been carried out for the first time in 1996 by Petr et al. during a single cyclic voltammetric experiment with the methyl-substituted p-phenylenediamine [71]. The experimental technique, including a special optical ESR cavity and an electrochemical cell for both ESR and UV–vis spectroscopy in transmission mode, was used. An improvement of UV–vis measurements was reached by the application of diode array spectrometers or spectrometers with CCD arrays in which a larger number of spectra can be recorded in the timescale of a single electrochemical experiment. The standard ESR optical transmission cavity was modified with respect to the construction of two optical openings. The front optical opening was adapted for the light guide of the continuous light source. The correct choice of the working electrode for studies of different kinds of compounds plays a crucial role. The use of a flat cell provides good conditions for sensitive ESR measurements but requires flat electrodes within the cell. The use of OTE electrodes in the flat cell compartment is often difficult and reduces the available potential range. Microstructured electrodes fulfill most requirements for a stable and sensitive behavior of the working electrode, but the penetration depth of the microwave field is too small for ESR measurements. It was found that a Pt-mesh electrode is the best experimental solution for doing first measurements of UV–vis–NIR spectra within the ESR cavity. Increasing the mesh size makes the ESR measurements more sensitive but leads to a broadening of the peaks in the cyclic voltammograms. An increase in the size of the holes in the Optical ESR cavity ESR spectrometer Potentiostat

CE RE WE

ESR console

Trigger signal Trigger signal Halogen lamp

Light guides

UV–vis detector NIR detector

Deuterium lamp

FIgURE 3.2 Simplified scheme of the triple in situ UV–vis–NIR/ESR spectroelectrochemical setup.

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platinum mesh or of the wire diameter was found to lead to a decrease in sensitivity and a delayed response of the UV measurements. A new kind of OTE based on very fine galvanically deposited metal meshes (1500 wires/in.), mechanically stabilized by insulating the edge and the connection wire by thermal laminating foil, was described for use in electrochemistry with common solvents by Neudeck and Kress [72]. The extremely fine mesh (with the extension of the diffusion layer exceeding the range of wire distance) with the well-insulated border enables a description of the voltammetric as well as the spectroscopic behavior with the simple model of finite planar diffusion. However, to get a signal proportional to the generated amount of product at the electrode, only very thin capillary slit cells with fast conversion times should be used. For conversion times in the range of several hundreds of milliseconds, the linearity is achieved after several seconds up to minutes, which allows to follow the decay of electrochemically generated radicals [72]. The electrochemical oxidation or reduction of molecules to more highly charged ions leads to the formation of differently charged compounds that can undergo comproportionation/disproportionation reactions in the bulk solution giving rise to products with different charges and spins. To overcome difficulties of the formation of several charged states in bulk at a distinct electrode potential, spectroelectrochemical measurements in thin-layer cells are recommended. It should be noted that even at thin-layer conditions, disproportionation of radical ions still takes place. At mesh electrodes, convection and edge effects can be observed for long electrolysis times, and an additional divergence appears at short times when the diffusion layers at each wire do not yet overlap. One can obtain an ESR signal proportional to the number of spins only if either spins are generated only inside the sensitive range of the cavity or the generation rate inside and outside the sensitive range is equal. By fixing these micromeshes in the center of a flat cell or a capillary slit cell, the electrochemical behavior can be simulated on the basis of planar finite diffusion and from the integration of concentration profiles the spectroscopic time dependence is available [72]. For in situ thin-layer UV–vis–NIR/ESR spectroelectrochemistry, a three-electrode arrangement with a laminated working electrode with a gold-μ-mesh, a platinum wire as a counter electrode, and a silver wire as a pseudoreference electrode was used [73]. To reach the thin-layer conditions, the electrolyte volume was reduced by inert foil sheets inserted into the flat cell as illustrated schematically in Figure 3.3. Commercially available thermal lamination foils of different thicknesses have been shown to be resistant against solvents such as acetonitrile, dimethyl sulfoxide, dimethylformamide, and others that are frequently used for electrochemical investigations. The lamination technique is applied to give micromeshes a sufficient mechanical stability. In this state, they can be handled as optical transparent electrodes with insulated edges. By thermal lamination of the micromesh between two laminating foils with holes for the active electrode surface, such meshes can be mechanically stabilized, insulated, and connected with a wire to have an insulated electrical contact to the working electrode. But the absorption spectra of diamagnetic structures are overlapping in mixtures of intermediates and/or products, and UV–vis–NIR spectroscopy generally suffers from fairly low resolution. Therefore, other spectroscopic methods resulting in more detailed structural information have to be applied, including FT-IR and NMR spectroelectrochemistry. In many cases in the electrochemical generation of radicals and their detection using ESR, it is often necessary to distinguish between paramagnetic species that are formed by an electron transfer process and those produced in a side or follow-up reaction, particularly in rapidly reacting redox systems. This problem could be solved by application of the phase-selective second harmonic AC voltammetry [74–76]. In the case of spectroelectrochemistry, only the absorbance of the primary intermediates varies in the same way as the perturbation signal. 2. LIgA Electrodes in Voltammetric and Spectroelectrochemical Studies The honeycomb lithographic-galvanic (LIGA) structures can be understood as a system of tubes that permit a fast electrochemical conversion time by using the structures in a capillary slit with a

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FIgURE 3.3 Scheme of (a) spectroelectrochemical cell, (b) detail of the central part of standard spectroelectrochemical cell (front and profile views), and (c) detail of the central part of thin-layer spectroelectrochemical cell (front and profile views): (1, quartz ESR flat cell; 2, laminated working electrode; 3, Ag wire; 4, Pt-wire counter electrode; 5, gold foil; 6, gold-μ-mesh; 7, laminating foil; 8, electroactive surface, not laminated; 9, inert and optically transparent foil sheets; 10, electrolyte volume in standard ESR spectroelectrochemical cell; 11, electrolyte volume in a thin-layer cell). (Adapted with permission from Matis, M. et al., J. Phys. Chem. B, 114, 4451. Copyright 2010 American Chemical Society.) 5 mm

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Epoxy glue PPy-layer on gold

FIgURE 3.4 Schematic diagrams of gold-LIGA electrode (a) and polypyrrole-LIGA electrode (b). (Reproduced with permission from Springer Science + Business Media: J. Anal. Chem., 367, 2000, 314–319, Dunsch, L., Neudeck, A., and Rapta, P., Copyright (2000) Springer.)

slit width of less than 25 µm (Figure 3.4a). But until now, these structures are rather expensive and not commercially available, and for their use in in  situ ESR spectroelectrochemistry, we need a special cell design. The advantages of LIGA fabricated microstructured honeycomb electrodes were demonstrated for spectroelectrochemical cells with respect to the response time (the time necessary to generate the product in a sufficient layer thickness close to the electrode to be detectable by UV–vis–NIR spectroscopy) and to the conversion of the redox system in solution under thin-layer

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conditions [77]. Transmission UV–vis–NIR spectroscopy for several electrochemical applications can be performed in a special spectroelectrochemical cell based on the LIGA electrode and the two quartz rods, forming the walls of the cell and conducting the light beam through the cell. They are limiting the diffusion layer at the structured part of the working LIGA electrode. These microstructured LIGA electrodes can be used as well-defined models of porous electrodes at which redox processes occur under finite diffusion conditions. Such electrodes have been successfully used in the voltammetric and spectroelectrochemical study of various redox systems in both aqueous and nonaqueous solutions [77–80]. Additionally, the possibility to fabricate the well-defined microstructures from various organic conducting polymers was demonstrated by the electrochemical deposition of polypyrrole in molded LIGA forms at high current densities in aqueous solutions [81] (Figure 3.4b). The studies were done in a spectroelectrochemical cell based on microstructured gold produced by the LIGA technique or on microstructured conducting polypyrrole layers as working electrode materials. These electrodes with well-defined hexagonal holes exhibit a large number of applications in the field of thin-layer CV and spectroelectrochemistry. Besides their application in voltammetric studies, they can be used as a new kind of OTEs with an aspect ratio, that is, the quotient of the width of the walls surrounding a honeycomb (or hexagonal hole) and the height (or thickness) of the structure of up to 20. The response behavior is advantageous without decreasing the active optical length as the small size of the honeycombs limits the diffusion layer thickness and results in a very fast conversion of the starting redox component into the product inside the hexagonal holes. The structure height of several hundreds of micrometers permits an extended optical length of the light beam passing the honeycomb holes where the electrochemical reaction takes place. Therefore, the response and the conversion time can be realized by using structures with a high aspect ratio under constant sensitivity of the spectroscopic measurement. The complete conversion of the redox system during the electrochemical experiment in the range of hundreds of milliseconds enables the study of complex redox reactions accompanied by follow-up reactions in the millisecond range. Additionally, the preparation of such microstructured electrodes using conducting polymers opens new ways to study the redox processes on the conducting polymer surface employing spectroelectrochemical techniques. 3. In Situ UV–Vis–NIR/ESR Spectroelectrochemistry of Conducting Polymers The use of several in situ spectroelectrochemical techniques results in the direct detection of both the paramagnetic and diamagnetic species formed during electrochemical doping of organic polymers. The simultaneous use of ESR and optical spectroscopies does allow us to differentiate the nature of the charge carriers electrogenerated during p-doping and n-doping to get the individual spectrum of each intermediate. A UV–vis–NIR/ESR spectroelectrochemical cell equipped with a laminated indium-tin oxide (ITO) working electrode [82] was used in the investigation of various organic substrates that are potential electron- or hole-transporting materials. The two ITO-coated glass plates were thermally laminated with solvent-resistant nonconductive lamination foil with two openings, 3 mm in diameter, to obtain the small electrochemically active surface with an insulated electrical contact (Figure 3.5). The nonelectrically contacted ITO piece enables the simple measurement of reference spectra. By changing the position of the electrode up and down in the ESR cell, the reference or probe spectra were measured without any changes in the spectrometer setup. To control the exact position for the reference and probe openings, a black mask with a hole smaller than the active electrode surface was fixed onto the ESR cell. To avoid oscillations of the cell voltage, two counter electrodes can be used where the second counter electrode, connected by a resistor with the potentiostat, is situated above the flat part of the cell [71]. The measurement of the reference UV–vis–NIR spectra direct in the ESR cavity was possible using a specially constructed noncontacted ITO plate in the spectroelectrochemical cell. The simultaneous in situ UV–vis–NIR/ESR spectroelectrochemical technique enables the monitoring

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WE

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(b)

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FIgURE 3.5 Schematic diagrams of UV–vis–NIR/ESR spectroelectrochemical cell equipped with a laminated indium-tin oxide (ITO) working electrode. (a) Laminated ITO electrode with the nonelectrically contacted ITO piece enabling the simple measurement of reference spectra. (b) Position of the cell for the measurement of reference (left) and polymer (right) optical spectra (WE, working electrode, RE, reference electrode, CE1 and CE2, counter electrodes). (c) Real view of the polymer deposited on laminated ITO positioned in the ESR flat cell. (1, polymer deposited on electroactive part of the ITO electrode; 2, laminated ITO plate with a gold foil contact; 3, Ag wire pseudo-reference electrode; 4, ESR flat quartz cell.)

of both ESR silent and paramagnetic species present in the investigated system during redox cycling. The obtained sets of spectral and electrochemical data can help to find a plausible redox mechanism of the investigated solid redox-active substrates on the electrode as well as to separate the superimposed UV–vis–NIR spectra. Using the calibrated manganese ESR standard, the quantitative time dependences of the three oxidation states in the polymer layer were obtained during redox cycling. The quantitative cyclic voltammetric and ESR data can be used to separate the superimposed UV–vis–NIR spectra into those of the individual redox states using the leastsquares method [83]. 4. In Situ UV–Vis–NIR/ESR Spectroelectrochemistry at Variable Temperatures The triple in situ experimental technique for the simultaneous measurement of UV–vis–NIR and ESR spectra in electrochemical experiments at different temperatures, including a special optical ESR cavity and Bruker temperature control system, was published by Rapta and Dunsch  [84]. With the possibility of recording spectra at several temperatures and concentrations during oxidation or reduction, it was possible to favor the formation of the radical ion or the corresponding dimer. The quantitative in situ UV–vis/ESR spectroelectrochemical technique was extended to the NIR region to enlarge the spectroscopic characterization of reaction products. The spectroelectrochemical cell was modified for measurements at different temperatures in a standard ESR Dewar insert positioned in an optical ESR cavity. For the spectroelectrochemical ESR flat cell applied in the in  situ UV–vis–NIR/ESR cyclic voltammetric experiments at different temperatures, a laminated platinum-mesh electrode was used as a working electrode (Figure 3.6a).

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FIgURE 3.6 (a) Scheme of the low-temperature spectroelectrochemical cell: WE—laminated platinummesh working electrode; RE—silver wire pseudo-reference electrode; CE—platinum wire counter electrode; T—flexible teflon tube; Mn—glass capillary filled with magnesium oxide containing traces of Mn2+. (b) Vis– NIR spectra observed during cyclic voltammetry of fullerene C120 in o-DCB + 0.1 M TBABF4 at 298 K and (c) at 260 K (inset: schematic structure of C120 monoanion and representative ESR spectrum simultaneously measured during the cyclic voltammetric scan at first reduction peak; *two ESR lines of Mn(II)-ESR standard sample). (Reprinted from J. Electroanal. Chem., 507, Rapta, P. and Dunsch, L., 287–292, Copyright (2001), with permission from Elsevier.)

A silver wire served as a pseudo-reference electrode and a platinum wire as a counter electrode. A flexible Teflon tube positioned on the bottom of the cell enabled all bubbles to escape from the cell before the measurement. During the triple in situ experiment, the spectroelectrochemical cell was positioned in a standard insert Dewar located inside the ESR optical resonator. The standard variable temperature setup was used to control the cell temperature. The triple in situ spectroelectrochemical measurements were controlled by triggering the UV–vis–NIR and ESR spectrometers. As an example of the valuable use of such a setup, it was shown that the stability of the C120 dimer is much lower as compared to C120O as the C120 molecule readily undergoes dissociation upon a one-electron reduction. In the spectroelectrochemical response of C120 at room temperature, the C60 − • anion radicals dominate as the reaction product as illustrated in Figure 3.6b. Thus, the characteristic NIR band of negatively charged C120 − • at 1020 nm during the cathodic reduction at 260 K was observed (Figure 3.6c). The presence of a narrow ESR line points to very similar ESR characteristics of both C120 and C120O dimers and corresponds to the single ESR line of R2C60 − • fullerides.

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II.

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SPECTROELECTROCHEMISTRy OF ORgANIC COMPOUNDS IN SOLUTION OR AS SOLID LAyERS ON THE ELECTRODE SURFACE

More illustrative examples for spectrochemical methods discussed earlier are given in the next paragraphs with examples of electrode reactions both in solution and in the solid state.

A.

EMPTY FULLERENES

Fullerenes are good electron acceptors and undergo multiple reversible single-electron reductions in solution. For instance, electrochemical studies of C60 showed that it is able to accept up to six electrons [85], and similar behavior is also found for other fullerenes (see also Chapter 21). Many spectroelectrochemical studies of the reduction process were reported since the early 1990s [86]. The anionic states of fullerenes exhibit characteristic absorptions in the NIR range originating from LUMO → LUMO+N excitations (here LUMO is the lowest-unoccupied molecular orbital of the pristine fullerene, which is populated upon reduction) [87]. Such excitations are absent in the pristine fullerene, which makes absorption spectroscopy an especially convenient tool for spectroelectrochemical studies. In the ESR spectra, anion radicals of fullerenes exhibit single-line signals whose line width may be rather large due to the Jahn–Teller effect when the degenerate LUMO is partially filled. For example, the ESR line width of C60 − • at room temperature is approximately 40 Gauss (G) [88]. For fullerenes with lower symmetry, the typical line width is less than 1 G (e.g., 0.15 G for C82− • and 0.2–1.0 G for isomers of C84− •) [89–92]. For anion radicals with sharp ESR signals, a 13C satellite structure can also be observed and gives information on the spin density distribution. The g-factor of monoanion radicals of fullerenes, g = 2.001–2.002, is close to the free-electron value (2.0023). A special exception is C60 − •, whose g-factor, 1.999–2.000, is noticeably smaller than in all other fullerenes because of the unquenched orbital momentum [88]. The g-factors of the trianions are usually shifted by approximately 0.001 to larger values than in corresponding monoanions. Especially beneficial can be a combination of ESR and absorption spectroscopies in spectroelectrochemical studies as shown in the studies of C60 [93], C82 [90,92], and four isomers of C84 [89]. A spectroelectrochemical study of C60 revealed the problem of a spike signal of unknown origin often observed in the ESR study of C60 − • anion radicals [86]. The in situ UV–vis–NIR/ESR study showed that only the broad signal of C60 − • is found if the fullerene is reduced solely at the first reduction step. However, if the potential range in the voltammetric study is extended further into the cathodic range (to the second reduction or beyond), the sharp ESR signal appears during the second voltammetric scan already in the region of the first reduction peak upon C60 reduction in o-dichlorobenzene, which is the most common solvent used in electrochemical studies of fullerenes. Thus, the ESR spectroelectrochemical study emphasized the high reactivity of the dianion (and higher charged anionic states of C60) and showed that a reaction with the solvent can take place already at the second reduction step. Note that the integrated intensity of the sharp signal was less than 1% of the main C60 − • signal, and hence the extent of the reaction at the timescale of the CV experiment is not sufficient to be detected by electrochemical methods.

B.

DERIVATIVES OF FULLERENES

The 13C hyperfine structure in the ESR spectra of the charged radicals of non-derivatized fullerenes is usually not sufficiently informative for a detailed analysis of the spin density distribution because of the low natural abundance of the 13C isotope. The situation can be changed when fullerene is exohedrally functionalized with substituents carrying magnetic nuclei such as 1H or 19F. Besides, derivatization also changes the π-system of a fullerene via saturation of some C-sp2, and hence the NIR absorption features of the anion can also be affected substantially. The vis–NIR/ESR spectroelectrochemical studies of alkyl [94–97] and perfluoroalkyl [98–101] derivatives of fullerenes showed a resolved hyperfine structure in many anion radicals and changes in their NIR absorption patterns.

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0.0

E vs. C700/–, V 0.4 0.8

E vs. C700/–, V

IESR 1.1 V

IESR

–0.4 0.0 0.4 0.8 1.2

1.1 V

0.1 V E vs. C700/– 1.1 V

(a)

–0.3 V E vs. C70 0/– 1.1 V

(b)

Sim. Exp.

2G

61 70 69 60

6 71 72

1000

(c)

C70(CF3)2

–3

q=0 q = –1 q = –2

–4 E, eV

ΔA

1090

C70

–5 –6

400 (d)

600

800

1000 λ, nm

1200

1400

1600 (e)

FIgURE 3.7 (See color insert.) (a) ESR spectra (only positive part is shown) measured in situ during cyclic voltammetry of C70(CF3)2 at the first reduction; (b) the same at the first and the second reduction peaks. The insets in (a) and (b) show cycling voltammograms measured in the same experiments (scan rate 10 mV s−1, 0.1  M TBABF4 in CH2Cl2 supporting electrolyte; Pt-mesh working electrode). (c) ESR spectrum of the C70(CF3)2− • radical anion. Insets: spin density distribution in C70(CF3)2− •, enhancement of the 13C satellite features and their simulation, and a fragment of the Schlegel diagram. In all simulations, the a(13C) values for C71, C72, C61, C70, C69, C6, and C60 (see Schlegel diagram for numbering) are 6.40, 4.46, 3.78, −3.66, 3.37, 2.83, and −2.67 G, respectively, which gave the best fit to the experimental spectrum. (d) Difference vis–NIR spectra measured in situ during electrochemical reduction of C70(CF3)2 at the first and second reduction peaks. (e) MO levels in C70 and C70(CF3)2, black arrows denote HOMO–LUMO gaps, gray arrows show schematically the excitations corresponding to the strong NIR bands in the absorption spectra of the anions and dianions. (Popov, A.A., Shustova, N.B., Boltalina, O.V., Strauss, S.H., and Dunsch, L.: ChemPhysChem. 2008. 9. 431–438. Copyright Wiley-VCH Verlag GmbH & Co. KGaA. Reproduced with permission.)

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To illustrate a typical spectroelectrochemical study of a fullerene derivative, Figure 3.7a,b shows ESR spectra measured during CV of C70(CF3)2 at the first reduction step [99]. The ESR spectrum of C70(CF3)2− • (Figure 3.7c) exhibited (1) a prominent 1:3:3:1 quartet with a hyperfine coupling constant (a(19F)) of 0.323 ± 0.004 G and (2) 13C satellites. This quartet structure indicates that only one of the CF3 groups is strongly coupled to the unpaired electron (i.e., only three equivalent F atoms are measurably coupled to the unpaired electron). DFT calculations performed to interpret the ESR spectra showed that the hfc constant of the CF3 group with the C71 carbon atom (see Figure 3.7c) is indeed several times larger than the value predicted for the second CF3 group (with carbon atom C72), and hence the latter cannot be seen in the experimental spectrum. The well-resolved 13C hyperfine structure with a(13C) values up to 6.4 G was also interpreted based on the calculations and showed that the largest values correspond to the 13C atom in CF3 groups followed by several carbon atoms of the fullerene cage. The evolution of the difference vis–NIR spectrum of C70(CF3)2 measured in  situ during CV at the first and second reduction peaks is shown in Figure 3.7d. At the first reduction step, an absorption maximum at 1090  nm (1.138 eV) has appeared, while the new intense band at 1000 nm (1.240 eV) has developed at the second reduction step. The main NIR bands of C70(CF3)2 anions are blue-shifted vs. the analogous transitions in C70 anions, 1370 nm (0.905 eV) in C70 − • and 1170 nm (1.060 eV) in C702− [86]. The blue-shift vs. C70 − • NIR band was also observed by Kadish et  al. [97] for the C70(C6H5CH2)2− • anion radical, which exhibited two NIR absorption bands at 1062 and 1250 nm. In agreement with experimental data, time-dependent DFT computations of the electronic excitations of C70(CF3)2− • and C70(CF3)22 predicted one intense NIR excitation at 1.065 eV (oscillator strength f = 0.013) for the monoanion, which is blue-shifted to higher energy in the dianion (1.153 eV, f = 0.022). Analogous transitions in C70 − • and C702 are predicted at 0.872 ( f = 0.017) and 1.018 eV ( f = 0.036), respectively. On the basis of these computations, intense NIR bands in C70(CF3)2− • and C70 − • are assigned to SOMO → LUMO+3 and SOMO → LUMO+4 excitations, respectively (Figure 3.7e), which become HOMO → LUMO+2 and HOMO → LUMO+3 excitations in the respective dianions.

C.

ENDOHEDRAL FULLERENES

Endohedral fullerenes (see also Chapter 21) encapsulate ions, molecules, or clusters in the interior space of the carbon cage [102], which results in a more complex redox behavior in comparison to empty fullerenes [103]. Whereas the carbon cage is the only redox-active center in the latter, both fullerene cage and endohedral species can exhibit redox activity in endohedral fullerenes. Electrochemical methods alone can hardly distinguish cage- or cluster-based redox processes, but this problem can be efficiently addressed by ESR spectroelectrochemistry if encapsulated species have nuclei with a nonzero spin (such as Sc with nuclear spin I = 7/2). The hyperfine structure in the ESR spectra can reveal if the spin density in the cation or anion radicals is localized on the endohedral cluster (large coupling constants and a pronounced shift of the g-factor from the freeelectron value will then be observed) or on the carbon cage. Besides, unlike empty fullerenes, many endohedral cluster fullerenes exhibit electrochemically irreversible but chemically reversible reductions [104]. The nature of the follow-up chemical process is still not well understood, and in situ spectroelectrochemical studies can deliver valuable information on the reasons of such redox behavior. Thorough in situ UV–vis–NIR and ESR spectroelectrochemical studies were performed for Sc3N@C68, an endohedral fullerene with the scandium nitride cluster Sc3N encapsulated in the C68 carbon cage [105,106]. A CV study showed that at room temperature in o-dichlorobenzene solution, Sc3N@C68 has two reversible oxidation and two irreversible reductions. Figure 3.8a shows the CV curve measured at the first reduction step, whereas Figure 3.8b demonstrates the time dependence of the ESR and NIR absorption signals detected in situ during the measurement of the CV.

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0.8

g e

I, μA

f

Intensity of ESR signal

b 0.4 a

0.2 0.0 –0.2 –0.4 –0.6

3 mV s–1 –0.2

0.0

(a)

0.2

0.4

0.6

Absorbance at 1226 nm

d

0.6

Forward scan Back scan –0.23 V –0.23 V 0.77 V

0.8

E, V vs. Fc/Fc+

(b)

10 G

a(45Sc) = 1.28 G

(c)

(d)

Absorbance

g f e d c b a

SOMO HOMO(–1) HOMO(–2)

400

600

800

(e)

1000 1200 λ, nm

1400

HOMO(–3)

1600

(f )

FIgURE 3.8 (See color insert.) (a) Cyclic voltammogram of Sc3N@C68 at the first oxidation step; the letters denote times when ESR and absorption spectra were measured. (b) Normalized intensity of ESR and NIR absorption at 1226  nm. (c) ESR spectrum of Sc3N@C68+ •; (d) DFT-computed spin density in Sc3N@C68+ •; (e)  Vis–NIR absorption spectra measured during oxidation of Sc3N@C68. (f) MO energy levels in Sc3N@ C68 and intense NIR excitations in the cation (gray arrows). (Yang, S.F., Rapta, P., and Dunsch, L., Chem. Commun., 189, 2007. Reproduced by permission from The Royal Society of Chemistry. Copyright 2006.)

Figure 3.8c shows the ESR spectrum measured during electrochemical oxidation of Sc3N@C68 at the first oxidation step. The hyperfine structure of the ESR signal comprising 22 equidistant lines is consistent with three equivalent Sc atoms. The small value of the hfc constant, 1.28 G, shows that the Sc atoms are weakly coupled to the unpaired spin in the cation radical. In good agreement with the experimental data, DFT calculations showed that the spin density in Sc3N@C68+ • is

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predominantly localized on the carbon cage (Figure 3.8d). Vis–NIR absorption spectra measured simultaneously with ESR showed that formation of the cation radical is accompanied by the two NIR absorption bands at 1097 and 1220 nm (Figure 3.8e). With the help of TD-DFT calculations, these absorptions were assigned to the HOMO–2 → SOMO and HOMO−3 → SOMO excitations, respectively (Figure 3.8f). Coincidence of the normalized ESR and absorption intensity time profiles obtained in the spectroelectrochemical measurements (Figure 3.8b) proved that the paramagnetic Sc3N@C68+ • cation radical is the only product of the electrochemical process. Formation of the paramagnetic anion radical with a 45Sc hfc value of 1.75 G was also detected by ESR at the first reduction step [105]. However, the integral intensity of the signal was approximately 1 order of magnitude lower than that of the cation measured under the same conditions. This result clearly shows that the main product of the electrochemical reduction is diamagnetic, which can be tentatively explained by the reversible formation of the diamagnetic single-bonded dianionic dimer: 2 Sc3N@C68− • ⇌ (Sc3N@C68)2−. Whereas Sc3N@C68 is an example of the endohedral metallofullerene with cage-based redox activity, Sc3N@C80 (which has the same endohedral cluster encapsulated in a different fullerene cage) is an example of an endohedral metallofullerene with cluster-based reduction as illustrated by ESR spectroscopic studies of its anion radical. The 45Sc hfc constant of Sc3N@C80 − • is 55.6 G (compared to 1.28 G in Sc3N@C68− •), and its g-factor is shifted to approximately 1.999 [107]. ESR spectroelectrochemical studies of exohedrally functionalized Sc3N@C80 derivatives, such as its monopyrrolidino-adduct and Sc3N@C80(CF3)2, showed that cluster-based reduction is still observed; however, hfc constants are smaller than in the anion radical of pristine Sc3N@C80 [108,109]. That is, the spin density in the anion radicals is partially transferred to the carbon cage, and the extent of this transfer is strongly controlled by the exohedral addition pattern. For Sc3N@ C80(CF3)2, ESR spectra were measured both at the first and third reduction steps and revealed substantially different hyperfine structure in monoanion (a(45Sc) = 2 × 9.34, 10.7 G) and trianion (a(45Sc) = 2 × 10.8, 49.2 G) [105]. The most pronounced effect of the spin density distribution in the electrochemically generated ion radicals was found in the recently studied Sc4O2@C80 with a mixed-valence state of the Sc atoms (Figure 3.9a) [110]. The compound exhibited two reversible redox steps in both the cathodic and anodic ranges, and DFT calculations showed that the HOMO and the LUMO are localized on the Sc4O2 cluster (Figure 3.9a,b). ESR spectra measured during the first reduction and oxidation steps of Sc4O2@C80 exhibited rich hyperfine structure corresponding to two pairs of nonequivalent Sc atoms (Figure 3.9c and d). Spin density is primarily localized on the Sc atoms either in anion or cation radicals, but the magnitude of the hfc constants is dramatically different. Whereas in the anion radical the a(45Sc) constants, 2 × 2.6 G and 2 × 27.4 G, are rather moderate, the cation exhibited very high hfc constants for one of the two pairs of Sc atoms: a(45Sc) = 2 × 19.0 G and 2 × 150.4 G. Such a large value was explained by the nature of the HOMO of Sc4O2@C80 having Sc–Sc bonding character with large 4s contributions (note that the isotropic hfc constant is proportional to the spin density at the nuclei and is therefore determined by the spin contribution from the s orbital since orbitals with higher angular momenta have nodes at nuclei). Thus, the ESR spectroelectrochemical study of Sc4O2@C80 not only proved the redox activity of the endohedral cluster but also provided information on metal–metal bonding in the molecule.

D. DIMERIzATION REACTIONS Our strong interest in the past was focused also to investigate the driving force of dimerization in monomer/oligomer structures. Monodisperse organic compounds with low molecular weight such as naphthalene diimides, thianthrenes, bipyrrols, bithiophenes, and extended conjugated oligomers, for example, oligothiophenes, oligophenylenes, and oligopyrroles were used for such studies (see Reference 111 and cited therein). For structural studies of the dimerization reaction products at electrode surfaces, the in situ UV–vis–NIR/ESR technique was applied [84,112–114].

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ScIII O

HOMO ScII

LUMO –2.5 –2.0 –1.5 –1.0 –0.5

(a)

0.0

0.5

1.0

E, V vs. Fe(Cp)20/+

(b)

* Exp.

Exp. 500 G

Sim.

2000 2400 2800 3200 3600 4000 4400 4800 Magnetic field, G (c)

50 G Sim.

3300 (d)

3500 3600 3400 Magnetic field, G

3700

FIgURE 3.9 (a) Frontier orbitals of the Sc4O2@C80 molecule and enlarged view on the Sc4O2 cluster showing two types of Sc. (b) Cyclic and square-wave voltammograms of Sc4O2@C80. (c) ESR spectrum of the Sc4O2@C80+ • cation; (d) ESR spectrum of the Sc4O2@C80 − • anion. An asterisk in (d) denotes an ESR signal of an unidentified impurity; intensity of the impurity signal is less than 1% of the total ESR intensity. (Reprinted with permission from Popov, A.A., Chen, N., Pinzón, J.R., Stevenson, S., Echegoyen, L.A., and Dunsch, L., Am. Chem. Soc., 134, 19607–19618. Copyright (2012) American Chemical Society.)

Oligothiophenes represent an attractive material for organic electronics as the type, number, and position of substituents, and the number of monomers in oligothiophenes can influence both the stabilization of charged states and the solubility of oligomers. The influence of the molecular structure on the stabilization of charged states was studied in detail by in situ UV–vis–NIR/ESR spectroelectrochemistry at a novel α,ω-dicyano substituted β,β′-dibutylquaterthiophene (DCNDBQT, Figure 3.10a) [115]. In this compound, the alkyl side chains result in sufficient solubility in organic solvents, while the electron acceptor cyano groups promote reduction processes. The CV of DCNDBQT points to a two-step oxidation (the first one is reversible, whereas the second one is irreversible at moderate scan rates) and a single reduction step (Figure 3.10b) [116]. The first oxidation and reduction processes should result in radical ions that are to be followed by in situ UV–vis–NIR/ESR spectroelectrochemical analysis. Direct evidence for the radical species of DCNDBQT was supplied by ESR spectroelectrochemistry (Figure 3.10c and d). The g-factor of the cation radical (g = 2.0025) is close to that of the free electron. The ESR spectrum with hyperfine structure indicates the predominant contribution of carbon atom orbitals to the SOMO of the cation radical with a slight contribution of nitrogen atoms. The higher g-factor of the anion radical (g = 2.0042) points to a higher portion of the sulfur atoms to the SOMO as compared to the cation radical.

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5 μA

In Situ Spectroelectrochemistry of Organic Compounds

Oxidation

CH3 Reduction S NC

S

CN

S

S

H3C

–2.0 –1.6 –1.2 –0.8 –0.4 0.0 0.4 0.8 1.2 1.6 2.0

(a)

E, V vs. DmFc/DmFc+

(b)

Experiment

Experiment

Simulation Simulation 0.5 mT 0.5 mT (d)

(c) 0.5

0.3

+

0.2

0.3 +

0.2 0.1 0.0

Rel. absorbance

Rel. absorbance

0.4 – –

0.1 0.0

–0.1 –0.1

–0.2 400 (e)

600

800 1000 1200 1400 1600 Wavelength, nm

400 (f )

600

800 1000 1200 1400 1600 Wavelength, nm

FIgURE 3.10 Structure of DCNDBQT (a), cyclic voltammetry (b), ESR spectra of cation (c) and anion (d)  radical, and UV–vis–NIR spectra of DCNDBQT recorded during the first oxidation (e) and reduction (f) steps. (Reprinted with permission from Haubner, K., Jaehne, E., Tarabek, J., Lukeš, V., and Dunsch, L., J. Phys. Chem. A, 114, 11545–11551. Copyright (2010) American Chemical Society.)

The quantitative ESR analysis of the first cathodic and anodic steps was done. Referring the charge transferred to the spin number of the cation radical of DCNDBQT, a ratio charge/ spin of 2.1 is found. That is, less than 50% of the charge units are used in the formation of a stable radical, while the higher percentage is forming a diamagnetic structure presumably by dimerization. The reversible formation of a diamagnetic π-dimer from two cation radicals was proven by in situ UV–vis–NIR and temperature-dependent in situ NMR spectroelectrochemical measurements [117]. Figure 3.10e and f shows the UV–vis–NIR spectra during the first electron transfer in the anodic and cathodic ranges. The experimental absorption maxima together with the calculated ones (by TD-DFT method) at the first anodic step prove the presence of a cation and an anion radical. The two bands at 646 and 1050 nm for DCNDBQT + • are assigned to SOMO → LUMO

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and HOMO → SOMO transitions, respectively, while for DCNDBQT − •, the bands at 752 and 1506 nm can be assigned to HOMO → SOMO and SOMO → LUMO transitions, respectively. The additional optical bands (554, 906, and 1294  nm for the cation radical and 492, 908, and 1260 nm for the anion radical) can be attributed to the formation of corresponding dimer structures. The optical spectra of the π-dimer present two optical π–π* transitions blue-shifted in relation to the optical bands of the corresponding radical ions and a third band at longer wavelengths (1294 and 1260 nm for the cation and anion radicals, respectively) assigned to charge transfer between the rings of the dimer. The reduction peak of DCNDBQT in the cathodic range results in even lower extent of an anion radical. The quantitative ESR spectroelectrochemistry showed that for the anion radical, the ratio charge/spin is as high as 3.0. The anion radical is therefore less stable at the experimental conditions used and transformed into dimeric and trimeric structures or even polymers as shown by ex situ MALDI-TOF mass spectrometry and indicated by in situ NMR spectroelectrochemistry [117]. The quantitative ESR spectroelectrochemistry supports the existence of follow-up reactions with the formation of diamagnetic structures for both the cation and anion radicals.

E.

CONDUCTING POLYMERS

The spectroelectrochemistry of solid structures at electrodes can be advantageously studied by in situ UV–vis–NIR/ESR spectroelectrochemistry to follow the influence of the chemical structure on the redox reactions of the polymer [118,119]. The chemical structure of polyaniline (PANI) is often given as a linear model with a linear arrangement of the monomers (Figure 3.11a). As the linear form of PANI is often applied in studies of charge transfer reactions, it is to be clarified whether NH

NH

NH

R

N NH

NH

N

R

NH

NH

HSO4–

(a) Reduced state ESR maximum After the first peak in CV After the second peak in CV

Absorption at 460 nm ESR Current

1310 620 Normalized absorption/current/ single integrated ESR signal

Absorbance

0.2

460

0.1

0.0 400

(b)

600

800

1000 1200 1400 1600 1800 Wavelength, nm

1.0

0.5

0.0

–0.6

(c)

–0.4

–0.2

0.0

0.2

0.4

0.6

Potential, V

FIgURE 3.11 Structure of PANI with phenazine branches (a), UV–vis–NIR spectra of the polymer recorded during its oxidation (b), and potential dependence of the current, ESR intensity, and absorption at 460 nm (c). (Reprinted with permission from Dmitrieva, E., Harima, Y., and Dunsch, L., J. Phys. Chem. B, 113, 16131– 16141. Copyright (2009) American Chemical Society.)

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only linear chains are existent in such structures of so-called “linear” PANI and whether their real structure has any influence on the stabilization of charged states in such kind of polymer. Especially the role of phenazine structures in the stabilization of charged states in PANI is of interest for the use of differently prepared PANI. The IR and ESR studies on “linear” PANI and in several copolymers of aniline and a phenazine derivative have provided evidence on the existence of phenazine units in the polymer structure [120]. However, the branching structure caused by the presence of the phenazine units in the polymer chains stabilizes different charged states in the polymer upon p-doping as shown by the combination of in situ UV–vis–NIR/ESR and ATR–FTIR spectroelectrochemistry. The detailed study of charged states in conducting polymers is of high importance because the existence of a polaron pair and that of a polaron can be differentiated by in  situ UV–vis–NIR/ ESR spectroelectrochemistry, and their electrode potential–dependent formation is available now to give new insights into the formation of these charged states in electrochemically prepared PANI. The ESR spectrum of PANI in the doped state has a narrow ESR signal with a line width of 0.5 G, while a polymer containing phenazine only gives no ESR signal upon oxidation or reduction [121]. Obviously, some extent of a linear chain is required to stabilize a polaron near the phenazine-like units. The potential dependence of the polaron formation in PANI points to a potential difference between the half-peak current and the ESR intensity (as the signal of the polaron) to be 70 mV (Figure 3.11b and c). Therefore, the formation of the polaron is potential-delayed. It is the rare case for PANI that the polaron is not the primarily formed charged state at the polymer chain, but at first, the polaron pair is formed by a two-electron transfer, which can dissociate into two polarons at higher potentials. The absorption band at 460 nm increases together with charge injection into the polymer and can be attributed to the polaron pair. The absorption peaks of the polaron might be hidden by the absorption of the polaron pairs, but the existence of the polaron can be clearly proven by the simultaneous ESR measurements. Due to the shift of the maximum of the polaron formation from the peak current, it is concluded that the polaron can be formed both by oxidation of the neutral polymer at higher potentials and by dissociation of the polaron pair at the lower potential range. The mechanism for the formation of the charged states in PANI upon oxidation includes an early step in the charge injection as the polaron pair formation. By in situ UV–vis–NIR/ESR spectroelectrochemistry, the formation of charged states in protonated and unprotonated structures of emeraldines was followed [122]. The initial stage of oxidation in emeraldine salt is preferably the polaron structure, while π-dimers and polaron pairs are formed by oxidation of emeraldine base. By quantitative in situ ESR measurements, the spin concentration for commercially available PANI structures was determined. The number of monomer units per spin was calculated to be 1 spin/100 aniline units in emeraldine salt. To conclude, the combination of different in situ spectroelectrochemical methods allows to establish the complete reaction mechanism of the formation of charged states in PANI upon oxidation.

REFERENCES 1. Bard, A. J.; Faulkner, L. R. Electrochemical Methods—Fundamentals and Applications, 2nd edn. Wiley: New York, 2001. 2. Bond, A. M. Broadening Electrochemical Horizons. University Press: Oxford, U.K., 2002. 3. Marken, F.; Neudeck, A.; Bond, A. M. Cyclic voltammetry. In Electroanalytical Methods: Guide to Experiments and Applications, 2nd edn. Scholz, F., Ed. Springer: Berlin, Germany, 2009, pp. 57–106. 4. Gosser, D. K. Cyclic Voltammetry: Simulation and Analysis of Reaction Mechanisms. Wiley-Interscience: New York, 1993. 5. Christensen, P. A.; Hamnett, A. Techniques and Mechanisms in Electrochemistry. Chapman and Hall: New York, 1994. 6. Bard, A. J., Stratmann, M., Schäfer, H. J., Eds. Encyclopedia of Electrochemistry. Organic Electrochemistry, Vol. 8. Wiley-VCH: Weinheim, Germany, 2004.

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4

Surface Techniques Mohamed M. Chehimi and Jean Pinson

CONTENTS I. Electrochemistry ................................................................................................................... 192 II. Infrared and Raman Spectroscopy ....................................................................................... 194 III. X-Ray Photoelectron Spectroscopy ...................................................................................... 197 IV. Time-of-Flight Secondary Ion Mass Spectroscopy .............................................................. 198 V. Atomic Force Microscopy ....................................................................................................200 VI. Miscellaneous ....................................................................................................................... 201 VII. Conclusion ............................................................................................................................202 References ......................................................................................................................................203 Chapter 42 describes the formation of surface-bound films by electrochemistry. In most cases, these films are nanometer thick, and their detection and characterization is by no way straightforward as it requests the use of high-performance surface-sensitive analytical tools. This chapter highlights a large set of methods that permit to interrogate the said thin films in terms of the following important characteristics: (1) the thickness of the film, (2) the chemical structure (Which chemical groups are present? What is the length of the oligomers?), (3) the compacity of the film (presence or absence of pinholes in the film; How many groups are bonded to the surface per cm2, Γsurf? How many groups are there in the whole thickness of the layer, Γvol?), (4) the presence of a chemical bond between the surface and the film, and (5) the final properties of the film (hydrophilic/hydrophobic surface, possible attachment of biomolecules, possible analytical measurements, permeability to molecular probes, etc.). During the design of experiments dealing with the modification of surfaces and the necessary ensuing analysis, some important points should be taken into account. In order to obtain a good characterization of the film, it is interesting to prepare the film with a functional group or a unique elemental marker that can be easily identified by one or several of the techniques described in the following; for example, nitroaromatic groups can be identified through their reversible system in aprotic medium (~−1.2 V vs. SCE), through their asymmetric and symmetric infrared (IR) bands (~1520 and 1340 cm−1), and through a characteristic x-ray photoelectron spectroscopy (XPS) signal at 406 eV clearly separated from other nitrogen signals; this is why this group has been widely used. Fluoro and perfluoro groups and long alkyl chains can also be easily characterized. Electrografting reactions, but also chemical or spontaneous grafting, often yield, besides the film on the surface, products in solution that can physisorb on the surface. Therefore, the surface must be carefully cleaned to be certain that only the grafted products are examined. This is generally done by rinsing the surface under sonication or in a Soxhlet extractor in good solvents of the expected side products. When examining thin films using different physicochemical techniques, one should be aware that depending on the method, different parts of the film are observed. For example, electrochemistry should, in principle, reflect the presence of all the electroactive groups in the film, but we will see in Section I, that this is not always the case. An IR beam penetrates the film on a distance of the order 1

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wavelength that is on a distance of the order of 1 µm, which for thin films amounts to the whole thickness of the film. In XPS, photoelectrons escape elastically from a depth of 5–10 nm, which means that for relatively thick films, only the outermost layers will be examined. In contrast, inelastically scattered electrons escape from very well-buried layers (up to 30 nm or more) and contribute to the shape and intensity of the survey spectral background. With time-of-flight secondary ion mass spectroscopy (ToF-SIMS), only the extreme surface of the film ( 0.2 (δ is a dimensionless distance from the electrode), more nodes are inserted when the layer relaxes into the solution (h-refinement). At the same time, the nodes move out from δ = 0 (bottom) into

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the solution (top). Furthermore, if the prescribed accuracy can be reached with fewer points, the grid is dynamically coarsened to improve efficiency of the solution process by deleting some nodes. Adaptive techniques were also applied in finite difference schemes [51,96,129,130] for a variety of problems with complex concentration profiles. Rosenbrock time stepping is applied and control is exerted over the solution itself and its gradient [129]. Recently, Amatore and coworkers [97,98] (see also the discussion in [131]) have proposed to use the rates (not rate constants!) of homogeneous chemical reactions as control criterion. This work has been extended to 2D simulations as well, including appropriate conformal mapping coordinate transformations [132]. These adaptive approaches are more complex than the basic algorithms described in Section II.C.2; however, they provide a flexible way to generate accurate simulation results even in problematic situations as discussed in Section II.C.3.

D.

SIMULATORS

The full advantage of electrochemical simulation comes with its general applicability to a wide variety of experiments and electrode reactions. Apart from writing program code that solves a particular, specialized set of model equations, it was always a goal of workers in this field to provide general tools for the use by not necessarily mathematically skilled or interested electrochemists. Such tools have been termed problem-solving environments [11,133,134], electrochemical simulators [10], or simulation packages [2]. Recent such programs try to achieve among others the following goals: • A large (virtually infinite) number of electrode reactions can be simulated, advantageously by providing a chemical notation to be supplied to the program. This requires automatic translation or compilation of the chemical equations describing the electrode reaction into the partial differential equation system with initial and boundary conditions and a sufficiently general solver algorithm. • The user of the program is required to invest only a minimum mathematical effort to generate the solution. Thus, application- rather than theory-oriented electrochemists are enabled to use simulation in their work. Automatic selection of algorithmic parameters that control the solution accuracy, however, is not always a simple task. • User input of model parameters such as rate constants should be tested for consistency. For example, in complex reaction networks, some rate or equilibrium parameters are correlated. This leads to thermodynamically superfluous reactions [60,135–137], and several electrochemical simulators do provide facilities to detect such a situation. However, stoichiometric consistency can also be checked [60,136], as well as charge consistency [60] if the chemical notation includes the specification of species charges. • A graphical user interface should provide choices to the user in a readily comprehensible format. The following, necessarily incomplete, list of electrochemical simulators (for other compilations, see [2,10,11]) does also show the historical development: • cvsim [24,138] was a FORTRAN-based program for cyclic voltammetric simulation in the EASI (electroanalytical simulation) package, which also included casim for chronoamperometric simulations) and gesim [31] (galvanostatic electrolysis simulation, including exhaustive and other variants). The programs only provided a limited set of electrode reaction mechanisms that were simulated with the orthogonal collocation technique. The user interface was text based. It should be noted that another program by the same name, but on the basis of finite differences, was propagated [139].

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• DigiSim [25] became the first commercialized electrochemical simulator [94,95,140,141]. It features a graphical user interface with display of the calculated curve, a mechanistic translator, semi-infinite and finite diffusion spaces, as well as hydrodynamic conditions. The temporal development of concentration profiles can be visualized (“the movie”), providing an impressive demonstration of the events occurring within the diffusion layer. Some limitations are: no adsorption processes and possibly inconsistent initial conditions (see Section II.C.3, item 3). DigiSim is based on finite difference algorithms. • DigiElch (presently in version 7 [142]) is also commercialized [143] and has been developed by one of the original DigiSim writers. It additionally includes—as compared to DigiSim—adsorption, more specific electron transfer laws, and concerted electron transfer steps (concerted proton–electron transfer, see Chapter 13). A limited trial version is available for test purposes. • ELSIM [133,134,144,145] has grown through various versions [146] and has a decidedly more complex user interface that, although requiring some mathematical background of the user, provides even higher flexibility. An extension of this system in the form of a “semiautomatic model builder for electroanalytical chemistry” (SAMBEAC) has been presented [147]. • ES-1 and ES-2 provide a simulation and visualization tool for assisted learning in the electrochemical context, in particular, finite diffusion conditions. The JAVA programs are integrated with HTML technology to form a dynamic textbook [148]. • EChem++ [13,29,30,60,85,149] uses still another approach different from the other simulators mentioned. It is an open-source project, providing free access [150] to the code base written in C++ [151]. In contrast to most of the other programs, it is developed on the Linux operating system. The design is based on object-oriented paradigms [149], which result in a rather flexible use of the program code. Besides having an electrochemical compiler [60] for the translation of chemical reaction sequences, EChem++ is characterized by a graphical user interface that flexibly allows the definition of some experimental details for both potential- and current-controlled experiments. For example, the user selects how the excitation of the electrode occurs, and complex combinations of steps, ramps, and other temporal functions of E (or i) are possible. Thus, based on the same core solver (adaptive finite elements), cyclic voltammetric [29,30,85], chronoamperometric (including multistep) [29,30], constant and programmed current chronopotentiometric [13], and other experiments can be simulated. Moreover, besides diffusional transport, electron transfer, and homogeneous kinetics (with separate definition of the reaction order and stoichiometry), adsorption, surface (heterogeneous) chemical reactions, and multielectrode arrangements (e.g., thin-layer, parallel-plane electrolysis) can be treated. Electron transfer modeled by Nernst conditions (equilibrium at x = 0 is attained at all times) and Butler–Volmer kinetics are available [30]. • KISSA-1D and KISSA-2D are based on the reaction rate control approach mentioned in Section II.C.4 [97,98,132,152]. (Hemi)spherical and (hemi)cylindrical geometries provide access to ultramicroelectrode simulation for general mechanisms. KISSA-1D was extended to the simulation of electrochemiluminescence experiments [153]. • Online services for electrochemical simulation were listed earlier (see [2], p. 279). However, at the time of this writing, only one of them [154] for a specific adsorption model was still available on the net. Electrochemical simulators are increasingly used in the electroorganic literature, and some examples of simulation results are also found in other chapters of this book (see, e.g., Figures 11.1, 11.4, and 11.5).

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APPLICATION

While Section II demonstrates the current state of simulation techniques, we will now discuss by means of examples what we can learn from applying such an approach in organic electrochemistry and how this can be done.

A.

SOLUTION QUALITY

We first describe how we judge the quality of simulation results. Is it possible to show that a particular simulation is true (verification, from Latin verus) or at least valid (validation; often by comparison to experimental data [155,156])? This is not a problem unique to electrochemical simulation, and at least from an epistemological point of view, the question has been answered no [155]. The best we can do is to make it highly probable that the model and the resulting simulation provide a good means to reproduce (calculated or experimental) data (confirmation [155]). Often, for electrochemical simulations, this is done by showing that solutions obtained at certain extreme parameter values replicate analytical results or other already accepted numerical values (limiting cases). For example, such benchmarking [155] was applied to spline orthogonal collocation simulations of cyclic voltammetric models based on reaction mechanisms with fast preceding or follow-up chemical equilibria coupled to an electron transfer [113]. The intrinsic problem of such a procedure is undoubtedly that the most interesting situations, where the model will have to be used later, are those that cannot be treated by limiting assumptions. However, an increasing number of such successful tests will add increasing credibility to the model. Some additional tests are as follows: confirming the correct operation of isolated algorithms within the simulation programs; refining the mesh formed by the nodes for discretization and checking convergence of the resulting simulation data (see the discussion in [2], p. 247ff., which also deals with consistency of the solution); analyzing the stability of the solution. Stability describes how an error at time t propagates while the solution develops further. This problem was already apparent in early finite difference simulations (Crank–Nicolson algorithm, leading to oscillatory behavior [93]) and has since been treated for a variety of related solution techniques [50,157–160]. Stability of electrochemical models has also been analyzed in strict mathematical terms (see, e.g., [161]). Owing to the inherent error control in adaptive schemes (see Section II.C.4), this seems to have become less important for such techniques.

B.

EXPLORATION AND PREDICTION

Simulations in organic electrochemistry are useful to explore the behavior of model reactions or experimental conditions that have not yet been (or even cannot at all be) applied to a real system. It can be quite instructive to predict, for example, the expected changes in observable i when adding or removing steps of the electrode reaction, varying rate and equilibrium constants, or using an unusual setup such as unsymmetrical potential–time triangles in cyclic voltammetry. Literature on this subject is extensive, and we discuss here only two selected recent examples. Fourier transform voltammetry, where a sinusoidal waveform is superimposed on the potential scan, has been advocated as an attractive alternative to classical (dc-) cyclic voltammetry [162]. The advantages of the technique were early demonstrated by simulation studies, and further “experimental strategies…based on knowledge gained from a perfectly general numerical simulation treatment…” were developed [163]. Behavioral patterns were observed in simulated data and expand the possibilities of dc voltammograms [164]. In particular, the fundamental, second, and higher harmonics [165] of the current response as well as its power spectrum (frequency dependence) were analyzed. Moreover, simulation studies provided estimates of experimental error effects, both random and systematic [166], as well as predictions of differences between Butler–Volmer and Marcus–Hush electron transfer kinetic models [167].

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Recently, a simulation study was performed to investigate systems in which the electrolyte is not or cannot be stirred between experiments. The calculations allow to estimate the effect of recovery of the initial diffusion layer during some waiting time in chronoamperometry [168], which would not easily be accessible by experiments. In a similar way, the development of concentrations close to the electrode during an initial period, where the electrode is held at its rest potential in the electrolyte, was simulated among others under natural convection conditions [49]. Again, these studies explore the behavior of a system that is barely subject to simple experimental investigation.

C. PRESENTATION OF SIMULATED RESULTS With the simplicity to quickly generate results of model calculations, it is no problem to produce large numbers of simulation data sets by systematically changing one or more parameters (see Section IV.A.5 in Chapter 1). For example, within a certain reaction mechanism, a particular rate constant k could be varied and the cyclic voltammetric response be calculated for each value of k. Often such work is done in the framework of dimensionless parameters. Given simple and fast simulation, it rather becomes difficult to deal with the quantity of numerical data. Working curves are a classical method to visualize simulation results, where we select a particular striking feature in the voltammogram (Figure 5.1), for example, the peak potential, and plot its value as a function of the dimensionless parameter characteristic for the electrode reaction (see, e.g., Figure 5.3a). Early examples of this technique historically mark the onset of cyclic voltammetric success [17,53,169]. While the curve is often presented only graphically, also mathematical functions may be provided [170–174]. The working curve approach has been extended to cases where two parameters determine the feature value, for example, the dependence of the peak potential difference ΔEp in cyclic voltammograms on rate constant and transfer coefficient of a quasireversible electron transfer [172,173] or the peak potential in an ECE mechanism variant (Figure 5.3b) on rate constants. Of course, here, the mathematical descriptions are highly advantageous. If more than two parameters are important, their combination into dimensionless quantities (see, e.g., Section IV.A.5 in Chapter 1) might be an effective way to reduce the complexity. This is particularly popular in the case of zone diagrams that are also a compact representation of the information content of a large number of simulations. Section III.B in Chapter 10 depicts some examples, as does Reference 175. A recent monograph [176] makes extensive use of such graphical representations. In a zone diagram (Figure 5.3c), we show the behavior of the simulation results in terms of limiting cases. Within a zone, a particular feature, based, for example, on voltammetric peak current, varies only within an experimentally achievable limit of, for example, 5% and is assumed to change only insignificantly [177], or follows a characteristic dependency. Each zone then defines characteristic behavior. In addition, these diagrams help to assess effects exerted if a reaction or experimental parameter changes. Recently, a mapping technique (high-dimensional model representation) has been established [179], where a set of voltammetric curves and its dependencies on parameters are described by an expansion into functions that have fewer variables. Then, in particular, data storage needs are drastically decreased, while full voltammograms can be quickly regenerated.

D. QUALITATIVE AND QUANTITATIVE ANALYSIS OF MECHANISMS Simulations support the elucidation of electroorganic reaction mechanisms in a qualitative way. In general, a mechanistic hypothesis can be made likely by rejecting alternatives if the behavior of the latter does not comply with the experimental data (qualitative analysis) [180]. On the other hand, it is impossible to prove a certain mechanism to be unique as an explanation for given experimental

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Organic Electrochemistry 40 40

0

E p–E (mV)

0

–20 –40

I

EIp–E 0 (mV)

20

–60 –80

–160

–100 –4

–6

–2

2

0

4

log k

–1

6

log k1

(a)

6

6

log k 1

–6 –6

(b) c0

log ρ Disp 1

k1

Cl

0

DMSO

k2

δ

–1 Br

ECE

DMSO

–2

I DMSO

–3

Cl

CN ACN

DMSO

H-atom

Br Cl

–4

I

DMSO Br

–5

–4

–3

–2

–1

CN ACN CN ACN

DMSO

C

1

log σ 2

3

(c)

FIgURE 5.3 Data presentation by working curves and zone diagrams: (a) electrochemically initiated dimerization, 2D working curve for one parameter, half-peak potential Ep/2 (referred to formal potential E 0) as a function of model rate constant κ1 = k1c 0/τ with k1 = dimerization rate constant, c 0 = initial concentration of electroactive substrate, and τ = timescale parameter of experiment (=F/RTv for cyclic voltammetry); (b) ECE mechanism with reversible chemical step, 3D working surface for two parameters, peak potential of first peak EpI (referred to formal potential of first electron transfer E 0) as a function of model rate constants κ1 = k1/τ and κ−1 = k−1/τ (k1 and k−1 being the rate constants of the forward and reverse chemical reactions); and (c) zone diagram for concurrent reactions initiated by electrochemical reductive cleavage of aromatic halides in terms of competition parameters ρ and σ (logarithmic scales), the hatched zone defines mixed control. (Reprinted with permission from M’Halla, F., Pinson, J., and Savéant, J.M., J. Am. Chem. Soc., 102, 4120–4127. Copyright (1980) American Chemical Society.)

data, because there might be additional alternatives that have not been considered in comparison, but reproduce the data with similar or even better quality. A systematic procedure to add necessary reaction steps to a mechanism (model expansion) or to delete superfluous steps to comply with the principle of Ockham’s razor (see [181], but also note the critical remarks in [155]; model reduction) based on sensitivity analysis has been proposed in the electrochemical context [182]. The quantitative use of simulation is based on a likely reaction mechanism (i.e., a model of the electrode reaction): the main goal is then to determined kinetic, thermodynamic, and transport parameters pertinent to the real system. This is often called an inverse problem (already mentioned in Section II.C), because we intend to find the parameters that—used in the simulation—give a best

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fit to the experimental data. The inverse problem in the electrochemical context is complicated by the fact that in general we deal with a multiparameter problem, where we need to estimate several parameters simultaneously, the relation between experimental data and parameters is often nonlinear (see, e.g., the working curves in Figure 5.3), the parameters might be strongly correlated (e.g., rate and equilibrium constants), and a suitable mechanistic hypothesis must be available. As a basis, experimental data must be used that are free of artifacts. In particular, iR-drop (uncompensated resistance, see Section II.E.2 in Chapter 2) must be avoided, and background currents must have been subtracted. Comparison of experimental and simulated curves is done in various ways: • Data transformation, often linearization, for example, by semi-integration [183]. In a similar way, global analysis has been applied [184]. • Feature analysis, which relies on working curves or surfaces as discussed in Section III.C and concentrates on particularly noticeable points in the data. As a particular application, a multiparameter estimation approach based on mathematically described working surfaces has been discussed [172,173,185,186]. Nonlinear optimization techniques determine those parameter values that provide a best representation of the features’ changes as a function of experimental conditions. • Analysis of full curves [25,187]: This is sometimes termed fitting [140] and is often directly coupled to simulation (e.g., in DigiSim [140] or the “professional” version of DigiElch [142]). The difference between experimental and simulated curves, expressed as a sum of squared residuals of data points R, is guiding the estimation of optimal parameters, starting from an initial guess. Unfortunately, most optimization algorithms are subject to be caught in local minima of R, while the best fit will be provided by a unique global minimum of this quantity. The use of various starting parameter combinations [140] is mandatory and the results must be critically compared. It is also important to check that the results are physically meaningful [95], for example, rate constants must not be negative. The data must cover a large range of experimental parameters (concentration c, scan rate v, additives and reagents in the electrolyte). For example, a single voltammogram can often be simulated easily with good agreement. However, the same set of reaction parameters should be used for the simulation of a large number of current/potential curves recorded at several c and v to increase the credibility of the results (see the following example). Thus, fitting to a set of data rather than to a single curve is highly recommended. If the data values in such a set differ by a large factor (e.g., currents in cyclic voltammograms recorded at various v), a suitable scaling procedure must be employed to avoid incorrect weighting of residuals. Indeed, such algorithms are customary in simulators such as DigiSim (see Section II.D). The curves in Figure 5.4 are the results of fitting simulated to experimental data according to these ideas. Figure 5.4b shows a selection of curves for the oxidation of anisaldehyde phenylhydrazone [188,189] at various v and c that are compared to simulations with a single set of rate constants. The experimental cyclic voltammograms are background corrected and iR-drop compensation by positive feedback was used. The simulations were generated under the assumption of the reaction mechanism shown in Figure 5.4a and the reaction parameters given in the figure legend. The radical cation produced in the primary electron transfer dimerizes (second-order kinetics) and/or is deprotonated by another hydrazone molecule, and the neutral radical resulting from the second reaction is further oxidized (ECE sequence with two additional side reactions). The proton transfer between radical cation and neutral starting hydrazone was formulated in the model mechanism as a sequence of two equilibrium steps (first-order deprotonation of radical cation and second-order protonation of hydrazone). Several alternative mechanisms were tested and unable to even qualitatively reproduce the experimental curves: both radical cation dimerization and self-protonation steps are essential parts of the overall reaction.

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R1

R2

N N C H H +e– –e–

+

R1 R1

+

R2

N N C H H

NH N CH

and other products R1

NH N CH +

B B = R1

R2

Dimerization

N N C H H

R2

R2

HB+ N N C H

R2

+e– –e–

R1 (a)

FIgURE 5.4

+

N N C H

R2

Products

Anodic oxidation of anisaldehyde phenylhydrazone (R1 = H, R2 = OCH3) in dichloromethane/0.1 M NBu4PF6: (a) reaction mechanism.

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R1

Current, i · 106 (V)

1.6 1.1 0.6

10

40

8

30

6

20

4

10 0

2

–10

0

0.1

–20

–2 –0.4 0.15 0.20 0.25 0.30 0.35 0.40 0.45 0.50 0.55 0.60 –3.6 Current, i · 106 (V)

–3.1 –2.6 –2.1 –1.6 –1.1 –0.6 –0.1

· 106 (V)

Current, i

–0.4 0.15 0.20 0.25 0.30 0.35 0.40 0.45 0.50 0.55 0.60

0.15 0.20 0.25 0.30 0.35 0.40 0.45 0.50 0.55 0.60 16 14 12 10 8 6 4 2 0 –2 –4 0.15 0.20 0.25 0.30 0.35 0.40 0.45 0.50 0.55 0.60

7

31

6

26

5

21

4

16

3 11

2

6

1

1

0 –1 0.15 0.20 0.25 0.30 0.35 0.40 0.45 0.50 0.55 0.60

(b)

Potential, E (V)

–4 0.15 0.20 0.25 0.30 0.35 0.40 0.45 0.50 0.55 0.60 Potential, E (V)

–30 0.15 0.20 0.25 0.30 0.35 0.40 0.45 0.50 0.55 0.60

60 40 20 0 –20 –40 0.15 0.20 0.25 0.30 0.35 0.40 0.45 0.50 0.55 0.60 130 110 90 70 50 30 10 –10 –30 –50 –70 0.15 0.20 0.25 0.30 0.35 0.40 0.45 0.50 0.55 0.60 Potential, E (V)

221

FIgURE 5.4 (Continued) Anodic oxidation of anisaldehyde phenylhydrazone (R1 = H, R2 = OCH3) in dichloromethane/0.1 M NBu4PF6: (b) experimental (symbols) and simulated (lines) cyclic voltammograms in the potential range of first oxidation: top row, c = 0.12 mM; middle row, c = 0.22 mM; bottom row, c = 0.41 mM; left column, v = 0.05 V s−1; middle column, v = 1 V s−1; right column, v = 20 V s−1. Potentials vs. ferrocene reference. Optimal parameters [188]: formal potential of hydrazone oxidation E0′ = +0.39 V, electron transfer rate constant of hydrazone oxidation ks = 0.4 cm s−1, dimerization rate constant kdim = 2.4 × 104 L mol−1 s−1, rate constant of radical cation deprotonation k deprot = 6.6 s−1, deprotonation equilibrium constant Kdeprot = 4.6 × 10 −8, rate constant of hydrazone protonation kprot = 1.1 × 108 L mol−1 s−1, protonation equilibrium constant Kprot = 5.2 × 105; values of the dimerization equilibrium constant (1010), the formal potential (0 V) and electron transfer rate constant (1 cm s−1) of neutral radical oxidation, as well as the transfer coefficients (α = 0.5) for both electron transfers were assumed as constants for fitting purposes.

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Organic Electrochemistry

The experimental scan rate and the concentration were varied by factors of 400 and approximately 3.5, respectively. Still, all experimental curves are reproduced in shape and intensity of forward and reverse peaks by the simulations with a single set of rate and equilibrium constants [188]. The fit is not perfect. In particular, experimental currents between the oxidation peak and the switching potential are sometimes larger, while some experimental peak currents are slightly smaller than the prediction. However, the example provides an estimate of the fitting quality that can be obtained in the real world. An additional note of caution concerns the fact that cyclic voltammograms commonly do show relatively little structure (as compared to, e.g., NMR or IR spectra). Then, relevant information is rather contained in the full ensemble of curves that have been recorded under varying experimental conditions. Consequently, the example demonstrates that it is much more important to be able to fit all curves satisfyingly, rather than a single curve perfectly. This also minimizes the danger of using an inappropriate mechanistic hypothesis: Single voltammograms can often be reproduced by several mechanisms (with certain kinetic and transport parameters) and the determination of mechanism and parameters does not give a unique solution. This is much less likely if a large ensemble of curves can be fitted using the same set of parameters. If a reasonable fit between experiments and simulation is found only in some of the voltammograms, for example, only for high scan rates, this might indicate that the mechanism is inappropriate, for example, missing a chemical reaction that becomes important at slow scan rates.

IV. CONCLUSION Simulation provides a widely used and valuable tool for the study of electroorganic reactions. A large variety of numerical algorithms is discussed in the literature, and comprehensive computer programs are available for the organic electrochemist. These programs allow the exploration of experimental conditions and reaction mechanisms and their effect on electrochemical measurements. On the basis of artifact-free experimental data, simulations support the estimation of kinetic, thermodynamic, and transport parameters. Thus, redox-active organic compounds are characterized and basic information for the planning of organic electrosyntheses is provided. Note added in proof: While this Chapter was in the proof reading state, a comprehensive work on the use of integral equations (see, Section II.C.1) was published [190].

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6

Theoretical Calculation of Reduction Potentials Junming Ho, Michelle L. Coote, Christopher J. Cramer, and Donald G. Truhlar

CONTENTS I. II.

Introduction .......................................................................................................................... 229 Formal Definitions, Electrochemical Concepts, and Basic Considerations ......................... 231 A. Ionization Potentials and Electron Affinities ................................................................ 231 B. Standard versus Formal Potentials ............................................................................... 232 C. Cyclic Voltammetry ......................................................................................................233 D. Effects of Protonation ................................................................................................... 234 E. Reversible and Irreversible Redox Processes................................................................ 235 F. Liquid Junction Potentials............................................................................................. 235 G. Reference Electrodes .................................................................................................... 236 III. Computation of Reduction Potentials ................................................................................... 236 A. Gas-Phase Free Energies of Reaction ........................................................................... 237 1. Gibbs Free Energy and the Treatment of Nuclear Motion..................................... 237 2. Electronic Energies of Atoms and Molecules ........................................................ 240 3. Standard State of the Electron ............................................................................... 241 B. Free Energies of Solvation ............................................................................................ 242 1. Absolute Potential of the Aqueous SHE ................................................................ 243 2. Nonaqueous Systems .............................................................................................244 C. Standard States ............................................................................................................. 245 D. Rates of Electron Transfer ............................................................................................ 247 IV. Examples............................................................................................................................... 247 A. Aqueous Standard One-Electron Reduction Potentials of Nitroxides and Quinones... 247 B. Chemically Irreversible Processes—Reductive Dechlorination .................................. 250 C. Constructing a Pourbaix Diagram for the Two-Electron Reduction of o-Chloranil .... 252 V. Concluding Remarks ............................................................................................................ 255 Acknowledgments.......................................................................................................................... 255 References ...................................................................................................................................... 255

I. INTRODUCTION The reduction potential is a direct measure of the thermodynamic feasibility of an oxidation– reduction half reaction; and it is fundamentally important in many aspects of organic, bioinorganic, and environmental chemistry, as well as in biology and materials science. The design of rational strategies for tuning the redox properties of compounds depends on understanding the key molecular features that dictate the reduction potential. As an example, in environmental chemistry, chlorinated aliphatic compounds are common environmental contaminants due to their widespread use as solvents and degreasers and are known to degrade via a reductive

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dehalogenation [1,2]; the environmental persistence of these compounds has been found to correlate with their relative reduction potentials, and the computation and measurement of these quantities is therefore valuable for understanding structure–activity trends and the design of environmentally friendly derivatives of these compounds [1,3–8]. Similarly, in biochemistry, nitroxides are a class of kinetically stable free radicals that have been widely studied as potential antioxidants against reactive oxygen species, which can lead to tissue injury and even cell death; both oxidation and reduction processes involving nitroxides are biologically relevant [9–12], and the ability to predict the redox potentials of nitroxides with various substituents and those embedded in rings can help prioritize synthetic targets for potentially biologically relevant antioxidants [13,14]. Reduction potentials are most straightforwardly defined when associated with readily reversible equilibria; in such instances, they contain equivalent information to equilibrium constants or free energy changes for electrochemical half reactions. In practice, the high reactivity of many species (e.g., organic radicals) participating in electrochemical reactions or the irreversibility or mechanistic complexity of redox reactions can make the direct experimental measurement of a corresponding reduction potential difficult. For this reason, computational chemistry offers a valuable alternative to experiment for the characterization of redox reactions. The theoretical calculation of any thermochemical quantity, including free energies and therefore including reduction potentials, usually takes advantage of the Born–Oppenheimer separation of electronic and nuclear motion, which ultimately reduces the problem to three steps: (1) the calculation of molecular potential energy surfaces by electronic structure calculations; (2) the treatment of nuclear motion, for example, vibrations; and (3) statistical mechanical averaging over relevant configurations, conformations, or solvent structures. Step (3) is often carried out by classical statistical mechanics and step (2) by quasiharmonic methods, whereas step (1) generally requires more expensive quantum mechanical (QM) calculations, which can limit the accuracy of predictions if sufficiently large systems make the application of accurate QM models impractical. However, the relatively recent development of efficient quantum chemical algorithms and powerful computer architectures has facilitated the quantitatively useful study of many reactions. Because most redox processes of practical interest occur in condensed phases, the development of reliable solvation models has also been critical to progress, and both implicit and explicit solvent models are now available such that well-chosen combinations of theoretical models have the potential to be used to make quantitative predictions of electrochemical quantities like reduction potentials. Although this chapter is concerned with thermodynamics, the reader should keep in mind that reactivity and biological activity also depend on kinetics. While kinetics is often correlated with thermodynamic descriptors such as reduction potentials, it also includes other factors whose complete discussion is beyond the scope of this chapter. Nevertheless, we will mention kinetic effects in some places because they are relevant to interpreting measurements. There are several approaches to calculating a condensed-phase reduction potential, ranging from phenomenological or theoretically guided linear free energy relationships (LFERs) correlating reduction potentials with other computed (or experimental) observables to direct calculations of reduction potentials. When using LFERs, computed properties are again often obtained by QM electronic structure calculations. Calculated or measured properties that may be correlated with reduction potentials include ionization energies and electron affinities in the gas phase, as well as energies of the frontier molecular orbitals (e.g., the highest occupied molecular orbital), and these quantities may be regressed on solution-phase reduction potentials in order to develop a predictive equation [15–28]. LFERs are appealing because they allow for very rapid evaluation of reduction potentials, which is especially important, for example, in high-throughput screening of large databases of drug candidates. The implicit assumption of such an approach is that the errors associated with neglecting contributions to the reduction potential that do not correlate with the chosen independent variables are negligible, as are errors associated with the level of theory used to compute these variables. In practice, LFERs may work well if the compounds under consideration

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are sufficiently similar to those used in the regression. The semiempirical nature of this approach means that it may be difficult to estimate the errors associated with these models, particularly when they are applied on compounds outside of the training set. When one attempts to calculate the reduction potentials directly, without linear regression against simpler quantities, typically only the most active portion of the system, for example, the solute and perhaps the first-solvent shell, is treated explicitly by quantum mechanics. The rest of the system is treated by molecular mechanics (MM), classical electrostatics, or both (although occasionally the whole system is treated by explicit quantum mechanics). Combining quantum mechanics for a primary subsystem with MM for the rest of the system is labeled QM/MM, and if the MM subsystem is the solvent, it is an example of an explicit solvent method that requires molecular dynamics (MD) or Monte Carlo (MC) methods to ensemble average the solvent. MD and MC free energy simulations permit examination of solvent structure and reorganization [29–32]. Methods based on classical electrostatics usually replace the discrete solvent molecules by a dielectric continuum, so that the solvent and the ensemble average over solvent configurations both become implicit. QM/ MM and implicit-solvent treatments are the methods of choice for the study of redox potentials in condensed-phase and biological systems because treating the entire system quantum mechanically raises the cost so much that one is usually forced to use less reliable methods or to skimp on ensemble averaging. In this chapter, we will focus exclusively on methods based on thermodynamic cycles where solution-phase reduction Gibbs free energies are computed by combining gas-phase energetics with solvation free energies of the products and reactants. Such methods are also used extensively in solution-phase pKa predictions [33–35] as well as in studies of other condensed-phase reactions such as free-radical polymerization [36,37]. In the following, Section II presents some formal concepts in equilibrium electrochemical thermodynamics. The Section III is concerned with the implementation of the computational protocols. The Section IV presents some worked examples.

II.

FORMAL DEFINITIONS, ELECTROCHEMICAL CONCEPTS, AND bASIC CONSIDERATIONS

This section introduces some formal concepts in equilibrium electrochemical thermodynamics that are important for calculating solution-phase reduction potentials.

A.

IONIzATION POTENTIALS AND ELECTRON AFFINITIES

The adiabatic ionization energy, usually called the ionization potential, is the energy required to form a molecular or atomic cation in its ground state via the loss of an electron from the ground state of the neutral system in the gas phase. The vertical ionization energy applies to the change in electronic energy upon removal of an electron from the equilibrium structure of the neutral without change in geometry, again in the gas phase. For this reason, the two quantities are identical for an atom, and for a molecule the vertical ionization energy is almost always higher than its adiabatic counterpart. The electron affinity (EA) is defined similarly to the adiabatic ionization energy, and the vertical electron attachment energy is similar to the vertical ionization energy, but these quantities refer to minus the change in energy when a neutral system gains an electron. The adiabatic quantities correspond to enthalpy changes at 0 K: M(g) → M + (g) + e − , ∆H 0 = IP M(g) + e − → M − (g), ∆H 0 = −EA where the subscript denotes the temperature in units of kelvin.

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(6.1)

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STANDARD VERSUS FORMAL POTENTIALS

At the heart of electrochemical thermodynamics is the chemical potential (μ), which equals the molar Gibbs energy (G) for a pure substance and the partial molar Gibbs free energy for a component of a solution. For a species A in a solution,  γC  µ A = µA° + RT ln    = µ°A + RT ln(a) C 

(6.2)

where C is the concentration a small circle in a superscript denotes the value of a quantity in the standard state a and γ are the activity and activity coefficient, respectively The usual standard states in the gas phase are an ideal gas at a pressure of 1 atm or 1 bar (0.987 atm), and the usual standard states for solutes in liquid-phase solutions are ideal solutions at a concentration of 1 M (1 mol/L of solution) or 1 molal (1 mol of solute/kg of solvent). Notice that we have introduced the dimensionless activity coefficients γi defined by [38] ai = γ i

ci c°i

(6.2a)

If we apply Equation 6.2 to the reaction Ox + e − → Red

(6.3)

where Ox is the oxidant Red is the reductant the free molar energy of reaction is given by a  ∆G = ∆G° + RT ln Q = ∆G° + RT ln  Red   aOx 

(6.4)

where Q is the dimensionless reaction quotient. The relation between free energy and the maximum electrical work that can be performed, as expressed in terms of the electrode potential E of a half-cell [39], is ∆G = −nFE

(6.5)

where F is the Faraday constant (96,485 C mol−1) n is the number of electrons in the half reaction Combining this with Equation 6.4 yields the following Nernst equation [40]: E = E° +

RT  aOx  ln F  aRed 

(6.6)

where E° is the standard electrode potential, also called the standard-state potential or the half-cell potential.

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Notice that E equals E° when the activities of all species are 1. However, such standard-state conditions are often difficult to achieve in practice, and standard-state potentials are often replaced by formal potentials, E°′. Formal potentials are sometimes called conditional potentials to denote that they apply under specified conditions rather than under standard conditions [38]. Specifically, this quantity is the measured potential of the half-cell when the ratio of the total concentrations of oxidized and reduced species is unity and other specified substances (e.g., proton) are present at designated concentrations. For example, they can be defined to correspond to the half-cell potentials when the concentration quotients (Qc) in the Nernst equation equal 1: E = E °′ +

RT RT  COx  ln Qc = E °′ + ln F F  CRed 

(6.7)

Then the formal potential (E°′) is related to the standard reduction potential (E°) as follows: E °′ = E ° +

RT  γ Ox  ln   F  γ Red 

(6.8)

For example, the absolute potential of the normal hydrogen electrode is based on a concentration of the proton equal to 1 mol L−1 and is therefore a formal potential. This may be corrected to give the absolute potential of the standard hydrogen electrode (SHE) by taking into account the activity coefficient for a 1 mol L −1 solution of [H + ] in water, which has been estimated to be 0.8 [41]: E °SHE = ENHE −

RT ln ( γ H+ ) = ENHE + 0.006 V F

(6.9)

In this particular instance, activity effects account for only a small change (6 mV) in the potential [42]. As an example of a more extreme case, the formal potential of the Fe3+ /Fe2+ couple varies from 0.53 to 0.7 V in 10 and 1 mol L−1 HCl solutions, respectively [43]. Typically, experimental standard reduction potentials are obtained by assuming a functional form that models the dependence of the potential on ionic strength. A series of formal potential measurements is then carried out at different values of ionic strength, and they are extrapolated to zero ionic strength where the activity coefficients approach unity [43].

C.

CYCLIC VOLTAMMETRY

Cyclic voltammetry is commonly used in the determination of formal potentials, which may be extracted directly from a fully reversible cyclic voltammogram as the average (midpoint) of the anodic and cathodic peak potentials, Epa and Epc, or from the half-wave potential of a sigmoid curve in steady-state voltammetry [43], to give a half-wave or midpoint potential, E1/2. Because the measured half-wave potential is affected by diffusion (a nonthermodynamic effect), it is related to the formal potential by E1/ 2 =

Epc + Epa RT  DRed  = E °′ + ln 2 2nF  DOx 

(6.10)

where D Ox is the diffusion coefficient of Ox. When the diffusion coefficients of the oxidized and reduced species are very similar, the half-wave potential provides a good approximation to the formal potential.

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EFFECTS OF PROTONATION

In aqueous solution, thermodynamically favored proton transfer is usually rapid, and electrochemical measurements usually give reduction potentials for half reactions that include any thermodynamically favorable proton addition or loss. As such, an n-electron, m-proton half reaction can be represented in two possible ways: Ox + ne − → Ox n−

(6.11a)

Ox + ne − + mH + → H m Ox ( n−m )−

(6.11b)

with corresponding standard reduction potentials denoted E°(Ox/Ox n−) and E°(Ox, mH+/ H mOx(n−m)−), respectively. The potential for the latter equation is directly dependent on pH and is equal to the formal potential E°′(Ox, mH+/H mOx(n−m)−) when the concentrations of all species are 1 mol L −1: E = E °′ +

RT [Ox][H + ]m RT [Ox] RT ln = E °′ + ln − 2.303 nF [H m Ox] nF [H m Ox] F

m  n  pH  

(6.11c)

By monitoring how this cell potential varies with pH, it is possible to determine the electron– proton stoichiometry (m/n) of the electrochemical measurement. For example, consider quinones and their derivatives, which are electroactive organic compounds that play a vital role in a number of biochemical processes. These compounds can undergo either a two-electron reduction (Ox/Ox2−), a  two-electron–one-proton reduction (Ox, H+/HOx−), or a two-electron–two-proton reduction (Ox, 2H+/H2Ox), depending on the pH of the solution [44]. In Section IV.C, we illustrate how one constructs an E versus pH diagram, which is called a Pourbaix diagram [45–47] and is analogous to a chemical speciation plot or predominance zone diagram determined by pH. A measured formal potential is a good approximation to the standard reduction potential only when activity and kinetic effects associated with chemical reaction(s) are relatively minor. Where this is not the case, explicit treatment of these effects should be included in the calculations, or comparisons should be made with other experimental potentials that correspond more closely to infinite dilution and to thermodynamic control. Consider the reduction of nitroxide radicals in aqueous solution in Figure 6.1.

Net reaction

FIgURE 6.1 solution.

E(Ox/Ox–)

N O

+

e–

– N O

+

H+

N OH

+

H+

N O + 2H + e– +

1/K2

1/K1

E(Ox, 2H+/H2Ox)

– N O

N OH + OH N H + OH N H

The microspecies present in the one-electron reduction of a nitroxide radical in aqueous

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The measured half-wave potential E1/2 is related to the formal potential for the one-electron–twoproton (1e, 2H+) transfer reaction as follows [48,49]:

(

)

E1/ 2 = E °′ Ox, 2H + H 2Ox + +

RT ln K1K 2 + K1[H + ] + [H + ]2 F

(

)

(6.12)

where K1 and K2 are the equilibrium constants associated with the protonation steps. In previous work, it was found that explicit consideration of the prototropic equilibria was necessary to obtain good agreement with the experimental half-wave potentials [14]. In some cases, the experimental potential corresponds to that for a one-electron (1e) transfer E°′(Ox/Ox−), and this is related to the formal potential of the (1e, 2H+) reduction potential E°′(Ox, 2H+/H2Ox+) by

(

)

(

)

E °′ Ox Ox − = E °′ Ox, 2H + H 2Ox + +

RT ln( K1K 2 ) F

(6.13)

E. REVERSIBLE AND IRREVERSIBLE REDOX PROCESSES Occasionally, half-wave potentials are also reported for quasi-reversible cyclic voltammetry experiments with a back wave partially present; however, the reader should note that these are usually estimated values and therefore may not be well suited for quantitative comparisons. It is impossible to extract E1/2 from completely irreversible processes (no back wave) because of kinetic control of the current such that the Nernst equilibrium is established less quickly than the change in potential or because there are fast follow-up (side) reactions consuming the pertinent species. There are instances where the transfer of an electron to or from a neutral precursor leaves the resulting radical ion in an electronic ground state that is dissociative [1,50]. (The former process is called dissociative attachment [51], and the latter is called dissociative ionization [52].) Following the electron-transfer event, which is rapid on the time scale of nuclear motion, the ion relaxes along the dissociative coordinate, leading to the scission of one or more bonds. Typically, the energetics associated with this fragmentation are such that the electron-transfer event is effectively irreversible. Depending on whether the ion lives long enough to be reoxidized/rereduced on the return sweep, the back wave may be only partially present or completely absent in a cyclic voltammogram, in which case it is not possible to extract a half-wave potential. An example of such a chemically irreversible process is the reductive dehalogenation of haloalkanes. For such processes, the equilibrium potential may alternatively be defined as the Gibbs free energy associated with the overall process, which in this case is Cl Cl Cl

Cl Cl Cl

+ 2e– + H+

Cl Cl Cl

Cl Cl + Cl– H

(6.14)

F. LIQUID JUNCTION POTENTIALS The liquid junction potential arises whenever solutions with two different compositions come into contact. Its magnitude depends on the relative concentrations of the various ions at the boundary and on their relative mobilities. These potentials may be significant in cases where the solvent system changes across a junction (e.g., from acetonitrile or dimethyl formamide [DMF] on one side to aqueous on the other). The liquid junction potentials of a number of dissimilar solvent junctions have been determined to range from 10 to 200 mV, depending on the junction [53].

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G. REFERENCE ELECTRODES The conventional reference electrode for aqueous systems is the SHE, which has been assigned a potential of zero in experimental measurements. In theoretical calculations, the absolute (rather than relative) reduction potentials are often computed, and knowledge of the absolute potential of the SHE is essential for comparing computations with experiment. A schematic of a cell with the aqueous SHE as reference and an Ox/Red couple in solvent S is as follows: Pt | H + (aq) (aH+ = 1); H 2 (g) (pH2 = 1 atm) | Ox(S)(aOx ); Red(aRed ) | Pt

(6.15)

In this equation, the SHE is the anode (where oxidation takes place, on the left), and the Ox/Red couple is the cathode (where reduction takes place, on the right). The vertical lines indicate phase  is the standard boundaries. The cell voltage (ECathode − EAnode) is given by Equation 6.16 where EOx/Red potential of the Ox/Red couple (see Equation 6.6) and Ej is the liquid junction potential between the aqueous SHE and the solvent/electrolyte containing Ox and Red:    + E j = EOx/Red − ESHE + Ecell = EOx/Red − ESHE

RT  aOx ln F  aRed

  + Ej 

(6.16)

If all species are in their respective standard states, with the activity (or concentration, as an estimate for activity) equal to 1 mol L−1 for solutions and fugacity (or pressure, as an estimate for fugacity) equal to 1 bar for gases, then Equation 6.16 simplifies into   Ecell = EOx/Red − ESHE + Ej

(6.17)

Since the physical setup of an SHE is somewhat cumbersome, reduction potentials are often referenced to other electrodes. In laboratory measurements, a secondary reference electrode whose potential versus the SHE(aq) is well known is usually used. Examples include the (KCl) saturated calomel electrode (SCE) and the saturated silver/silver chloride electrode; the presence of saturated KCl in these electrodes leads to sharply reduced values of Ej. As such, in comparing with experiment, it is also important to examine the details of the experimental measurement to ascertain whether a correction for Ej is necessary in theoretical calculations. The conversion constants between different electrodes in aqueous solvents have been measured [54], and these may be used to convert reduction potentials that are referenced to SHE to other reference electrodes. For example, the potential of the SCE is 0.244 V relative to the SHE at 298 K in aqueous solution. Therefore, to convert values based on SHE to SCE, one needs to subtract 0.244 V.

III. COMPUTATION OF REDUCTION POTENTIALS As indicated in Equations 6.4 and 6.5, the standard-state Gibbs free energy change for a half reaction is the quantity required for computing a standard reduction potential. Since experimental reduction potentials are not measured in isolation but are instead measured relative to the potential of a reference electrode, theoretical calculations of reduction potentials are typically carried out either for a half-cell reaction (Figure 6.2, cycle A) with the subtraction of the reduction potential of the reference electrode (e.g., SHE) or on a full-cell reaction (Figure 6.2, cycle B). In Figure 6.2 we have introduced the general notation ∆GS for a standard-state free energy of solvation, which is the free energy change upon transfer from the gas phase (sometimes called air in the transfer literature) to the liquid solution.

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Theoretical Calculation of Reduction Potentials Cycle A M(aq) + –ΔGso(M)

ΔGrxn

e–(g)

ΔGso(M–)

ΔG = 0

+

M(g)

e–(g)

o ΔG gas

Cycle B M(aq) + 1/2H2(g) –ΔGso(M)

+ 1/2H2(g)

M–(g)

ΔG rxn

M–(aq) + H+(aq)

ΔGso(M–)

ΔG = 0

M(g)

M–(aq)

ΔG ogas

ΔGso(H+)

M–(g) + H+(g)

FIgURE 6.2 Thermodynamic cycles for calculating an absolute and relative reduction potential.

The corresponding reduction potentials are Ecell =

−∆Grxn (A) − ESHE nF

(cycle A)

(6.18)

and Ecell =

−∆Grxn (B) nF

(cycle B)

(6.19)

In principle, both cycles yield the same result. However, cycle B effectively uses calculated values of ESHE and ∆GS (H + ), whereas cycle A effectively uses empirical (accurate) values. Thus, cycle A is simpler, and in this cycle, the key ingredients for the calculation of a reduction potential are the gas-phase Gibbs free energy of reaction and the free energies of solvation of the reagents, that is, of the reactants and products.

A.

GAS-PHASE FREE ENERGIES OF REACTION

1. gibbs Free Energy and the Treatment of Nuclear Motion The Gibbs energy change of the gas-phase reaction shown in cycle A is simply the EA of M, EA(M), plus the thermal contribution to the Gibbs free energy (ΔG therm) of M− less that of M:  ∆Ggas = G  (M − ) − G  (M)

= U e (M − ) + ZPE(M − ) + ∆Gtherm (M − )  − U e (M) + ZPE(M) + ∆Gtherm (M)  = −EA(M) +  ∆Gtherm (M − ) − ∆Gtherm (M)  = −EA(M) + ∆∆Gtherm

(6.20)

where Ue denotes the Born–Oppenheimer equilibrium potential energy ZPE denotes the vibrational zero point energy ΔG therm denotes the thermal contribution to the free energy, that is, the part that vanishes at 0 K

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The thermal contribution includes the free energy due to multiple conformations (if present), rotations, and vibrational and electronic excitation. Note that the change in ZPE is included in the EA. We have neglected nuclear spin considerations, since the effect of nuclear spin cancels out in almost all cases, the main exception being the H2 molecule. It is useful to introduce the enthalpy at 0 K, which is labeled H0. Then H 0 = U e + ZPE

(6.21a)

 ∆Ggas = H 0 (M − ) + ∆Gtherm (M − ) − H 0 (M) − ∆Gtherm (M)

(6.21b)

and Equation 6.20 becomes

If the conformations, geometries, and vibrational frequencies of the charged molecule are very similar to those of the neutral and neither has low-lying electronically excited states, then the thermal correction to  the Gibbs energy of M− and M is likely to be similar and one could roughly estimate ∆Ggas as approximately equal to EA(M). In some cases however, the gain (or loss) of an electron can result in significant changes to the electronic structure of a molecule (e.g., quinones acquire an aromatic ring structure upon the gain of two electrons), and this approximation becomes unreliable. In such situations, the thermal corrections are sometimes calculated by assuming ideal gas behavior and the rigid-rotor harmonic oscilɶ from lator approximation, to arrive at analytic expressions for the molecular partition function (Q), which one can calculate the entropy (S), and the thermal contributions to the enthalpy (ΔHtherm) and the Gibbs free energy (ΔGtherm), which are evaluated from the following expressions: ɶ   ɶ + T  ∂ ln Q   S = R  ln Q   ∂T V   ɶ  ∂ ln Q ∆H therm = RT 2   + RT  ∂T V ∆Gtherm = ∆H therm − TS

(6.22)

ɶ is the molecular partition function with zero of energy at the ground state and the equivawhere Q ɶ is lent expression for the Gibbs free energy in terms of Q ɶ G = U e + ZPE + PV − RT ln Q

(6.23)

Furthermore, if one assumes that there is only one conformation and negligible coupling between electronic excitation, vibrations, and rotations, the molecular partition function can be separated into a product of partition functions associated with the translational, rotational, vibrational, and electronic motions: ɶ = qtransqrot qvibqelec Q

(6.24)

If we assume separability, the electronic partition function is ∞

qelec = ω1 +

 −ε i 

∑ ω exp  k T  i

i =2

(6.25)

B

where εi is the electronic energy (including nuclear repulsion but not vibrational energy) of level i ωi is the degeneracy of that level

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When the first electronic excitation is thermally inaccessible at room temperature, the electronic partition is well approximated by the degeneracy associated with the electronic ground state: (6.26)

qelec = ω1

For monatomic species, if the total electronic angular momentum associated with electronic state i is Ji, we have ωi = 2Ji + 1. For example, the ground state of a halogen atom is 2 P3/ 2 with J1 = 3/2, so ω1 = 4, and the first excited state is 2 P1/ 2 with J2 = 1/2 and ω2 = 2. Based on Equation 6.25 and the excitation energy of 0.109 eV of the first electronically excited state, the electronic partition function for chlorine atom at 298 K is therefore qelec = 4 + 2e −4.2 = 4.03

(6.27a)

Higher excited states make a negligible contribution in this case. The vibrational partition function is usually treated by the harmonic oscillator approximation or by a quasiharmonic approximation in which one uses the harmonic oscillator formulas but scales the frequencies [55–57] to account for anharmonicity (and for systematic deficiencies of the electronic structure method used to calculate the frequencies). The rotational partition function is usually treated classically. For molecules where there are multiple conformers that are close in energy to the lowest-energy structure, the conformational flexibility contributes to the G therm. If we again make a separable approximation, we can include this by putting another factor in Equation 6.24, yielding ɶ = qtransqrot qvibqelecqconf Q N conf

qconf =

 −∆U j  kBT 

∑ exp  j =1

(6.27b)

(6.27c)

where ΔUj is the potential energy difference of conformation j from the lowest one and the conformational partition function is summed over all the conformational space of the molecule, which is equivalent to performing a Boltzmann average over the Gibbs free energies of all the conformers. A much better approximation is to use Equation 6.24—or a more accurate analog with less separability approximations—to calculate a free energy Gj for each conformer. Then the free energy including all conformers is G = − RT ln

 −G j 

∑ exp  k T  j

(6.27d)

B

One should only include distinguishable conformers. However, even the number of distinguishable conformers grows rapidly with molecule size for chain molecules. For example, n-heptane has 59 distinguishable conformations [58]. Glucose has 2916 potential conformations [59]. Even the approximation of Equation 6.27d is far from realistic, though, if the barriers separating the conformers are not all high compared to k BT, both because the contributions of any one conformer are no longer independent and because the individual contributions differ from their harmonic values. If the barriers are low, the system must be treated as having one or more internal rotations. A theoretical formalism, based on internal coordinates and including intermode coupling, is available [60].

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In practice, a full conformational search typically involves at least 3N geometry optimizations where N is the number of rotatable bonds in the molecule that yield distinguishable structures (e.g., the C-3 to C-4 torsion in 1-butanol does not yield distinguishable structures). Additional considerations apply if one must consider ring isomerism as well as torsional isomerism. Therefore, a full conformational search is usually restricted to molecules with N ≤ 5. It is worth noting that a rough approximate upper bound on the effect of considering higher conformers is given by the case where there are Nconf conformers with energies, structures, and frequencies identical to those of the lowest-energy structure; then the error associated with not including the conformational partition function is RT ln(Nconf ). A variety of methods such as simulated annealing [61], MC methods [62], and an energy-directed tree search algorithm [63] have been developed for locating the lowest-energy conformer without having to sample the entire conformational space of the molecule. In principle, one should rank the conformers in terms of their Gibbs free energies as in Equation 6.23; however, this entails relatively expensive Hessian calculations, and in practice, the conformers are usually ranked in terms of their electronic energies (Ue). As a precaution, one could, at the end of the search, perform Hessian calculations only on conformers that are within some energy difference from the lowest-energy structure and rerank the conformers in terms of their Gibbs energies. The expressions for the partition functions as derived from the ideal gas, rigid-rotor harmonic oscillator approximation can be found in standard textbooks [64] and will not be presented here. A discussion of the potential sources of error in the application of these partition functions (e.g., breakdown of the harmonic oscillator approximation) and the errors that could arise from the assumptions used to derive them has been discussed elsewhere [65,66]. These treatments assume that the torsions are separable and may be identified with specific normal modes. When this is not the case, one must use the internal-coordinate nonseparable treatment mentioned earlier [60]. Having laid out the key ingredients for calculating a gas-phase Gibbs free energy, we now discuss possible levels of theory for calculating geometries, Born–Oppenheimer (electronic) energies, and free energies. Geometries are often calculated at lower levels of theory such as density functional theory (DFT) with a small basis set that can predict equilibrium geometries and vibrational frequencies (when scaled by appropriate scale factors [56,57,60]) reasonably well but is not usually sufficiently accurate for reaction energies. However, it is also usually possible to calculate geometries at the same level as reasonably reliable energies if one uses DFT with a modern density functional and a good basis set. 2. Electronic Energies of Atoms and Molecules Chemically accurate (errors of 5 kJ mol−1 or less) electronic energies of reaction can usually be achieved for small- and moderate-sized systems provided that electronic energies are calculated at high levels of theory, for example, CCSD(T) or QCISD(T), with very large one-electron basis sets incorporating high angular momentum basis functions. Here CC denotes coupled cluster theory, QCI denotes quadratic configuration interaction, SD denotes the inclusion of single and double excitations, and (T) denotes a quasiperturbative treatment of connected triple excitations [67]. One difficulty with electronic wave function theory (WFT) methods of this sort is the very slow convergence of the energies with respect to the size of the one-electron basis set. Furthermore, a CCSD(T) calculation formally scales as the seventh power of the number of atoms in the system [67] and is therefore restricted to relatively small molecular systems. Popular alternatives to large-basis-set CCSD(T) calculations are composite methods that have been designed to approximate high-level-correlated calculations using a series of lower cost calculations in conjunction with additivity and/or extrapolation routines. The Gaussian-n (e.g., G4 [68], G3 [69], G3(MP2) [70], and G3(MP2)-RAD [71,72]) methods with high-level corrections, multicoefficient correlation methods [73–82], the correlationconsistent composite approach [83,84], and CBS-X (e.g., where X is QB3 [85]) are examples of such methods. These methods involve some degree of empirical parameterization and are practical for medium-sized systems. By comparison, the Wn (n = 1–4) methods [86–89] have been designed to

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compute thermochemical properties with even higher accuracy (ca. 1 kJ mol−1), without empirical parameterization, but are also considerably more expensive and therefore limited to relatively small systems. For larger systems where even composite methods become computationally expensive, one could employ an ONIOM approximation [90,91] where the chemical system is partitioned into layers. The innermost layer is usually defined by the reaction center and its nearby substituents so that the chemistry of the reaction is modeled accurately. This layer is treated at the highest level of theory. The subsequent layer(s) are then treated at lower levels of theories. As an example, this approach has been successfully used to approximate the G3(MP2)-RAD calculations for a test set of 112 different radical reactions with a mean absolute deviation of 1.2 kJ mol−1 [92]. There are a large number of other shortcuts and “tricks of the trade,” for example, basis-set extrapolation [93–96] to ameliorate the aforementioned slow convergence, but these are too numerous to mention. 4 3 An important alternative to WFT is DFT. Here the computational work scales as N atom or N atom 7 rather than N atom, where Natom is the number of atoms in the system, but the accuracy depends on the quality of the exchange–correlation functional [97]. This quality is improving rapidly [98]. We next address relativistic effects, which begin to be energetically important at the level of chemical accuracy near the end of the first transition-metal series. There are two kinds of relativistic effects: (1) scalar relativistic effects and (2) spin–orbit coupling [99]. Scalar relativistic effects are most simply handled by replacing the core electrons with appropriate effective core potentials [100–105]; however, the accuracy can be low [106]. If all electrons are to be treated, the most rigorous approach makes use of the four-component Dirac spinor operator. More efficient approaches are based on two-component spinors; such methods can be derived from the four-component formulation through various transformations that lead either to the Douglas–Kroll–Hess Hamiltonian [107–109] or the zero-order-regular approximation [110]. Additional reduction to a one-component formulation yields the spin–orbit operator in its usual form and also a spin–spin interaction term [111]. Spin–orbit effects, associated with the coupling of spin and orbital angular momenta in a relativistic framework, are sometimes neglected in electronic structure calculations that make use of basis sets including relativistic pseudopotentials [99]. Rather, only scalar relativistic effects are included and computed energies represent averages over spin–orbit states, if they exist. Spin–orbit effects can be included through either perturbation theory or variational methods without sacrificing the simplicity of one-component computational models. When the relevant transition-metal compounds may be viewed as substantially ionic in character, a particularly simple approach is to estimate spin–orbit effects on standard reduction potentials by assuming the same spin–orbit coupling in the complexes as that for the bare ions, where the latter are usually available from experiment. 3. Standard State of the Electron In calculating ionization energies and electron attachment energies at nonzero temperatures or when calculating the free energy of reaction of processes like Equation 6.3, one needs to take into account the Gibbs free energy of the electron. There are two thermochemical conventions concerning the thermodynamics of the electron: (1) the electron convention (EC) and (2) the ion convention (IC). There are various literature reports giving slightly different calculated reduction potentials depending on which thermochemical convention of the electron is used. However, this should not be the case, and it originates from confusion regarding the definition of the zero of energy in the two conventions. An important point is that the Gibbs free energy obtained from a particular convention must be compatible with the quantum chemical calculation, that is, they need to have the same zero of energy. In quantum chemistry, it is convenient to define a zero of energy, at least temporarily, as corresponding to all nuclei and electrons being infinitely separated and at rest. With regard to the zero of energy for the free electron, the two conventions primarily differ in their definition of the formation enthalpy of the electron. The EC treats the reference state for electrons in the same way as for elements, that is, the enthalpy of formation is defined to be zero at all temperatures, ∆ f H T (e − ) = 0. On the other hand, the IC defines the standard enthalpy of formation of the electron to be equal to

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TAbLE 6.1 Thermodynamics of the Electron under the Various Thermochemical Conventionsa ∆ f H °298 ∆ f G°298 S °298 [ H °298 − H °0 ] G°298

EC–b

IC (bartmess)

IC–b

EC/IC–FD

0 0 20.979 6.197 −0.058

0 (0)b (0)b 0 (0)b

6.197 0 20.979 6.197 −0.058

0 0 22.734 3.146 −3.632

EC, electron convention; IC, ion convention; B, Boltzmann statistics; FD, Fermi–Dirac statistics. Enthalpies and free energies in kJ mol−1, entropies in J mol−1 K−1. b Defined values [114]. a

its integrated heat capacity at all temperatures [112,113]. Accordingly, under the two conventions, the enthalpy of formation of ions differs by the integrated heat capacity of the electron; the actual value depends on the statistical formalism used to treat the electron. Using Boltzmann statistics and the ideal gas model, the Gibbs energy of the electron is 0 kJ mol−1 at 298 K. However, since electrons are fermions, Fermi–Dirac statistics are more appropriate, and this yields a Gibbs energy of −3.6 kJ mol−1 at 298 K. Contrary to the earlier report by Bartmess [114], these values are the same under both conventions and the thermochemistry of the electron is summarized in Table 6.1. In the calculation of reduction potentials, it makes no difference which formalism or convention is used as long as these are used consistently for both the half-cell and the reference electrode.

B.

FREE ENERGIES OF SOLVATION

Continuum solvation models [115–118] have been designed to make accurate predictions of free energies of solvation. Free energies of solvation can then be combined with the gas-phase Gibbs energies in Equations 6.20 and 6.21 to obtain the Gibbs free energy of reaction in solution. In continuum solvation models, the solute is encapsulated in a molecular-shaped cavity embedded in a dielectric continuum. The solute is acted on by a reaction field, which is the field exerted on the solute by the polarized dielectric continuum, and the polarization of the solute by this field is calculated via the Poisson equation for a nonhomogeneous dielectric medium (the nonhomogeneous formulation [119] is required because ε is unity inside the cavity—because polarization is treated explicitly—but not unity outside the cavity where it is given the value of the solvent’s bulk dielectric constant). The reaction field is used to calculate the bulk-electrostatic contribution, which is then combined with the non-bulk-electrostatic terms to yield the solvation free energy. There are two contributions to the non–bulk electrostatics. One is the deviation of the true electrostatics from the electrostatics calculated using the bulk dielectric constant. The other is the nonelectrostatic portion of the solvation free energy. Some continuum solvent models such as the polarized continuum models (PCM, e.g., [IEF]-PCM [120] (IEF = integral equation formalism) and CPCM [121,122]) model the non-bulk-electrostatic and bulk-electrostatic terms independently; such models are called [123] type 3 models. Such models are less accurate than type 4 models [49,118,123–131], which are models in which the non-bulk-electrostatic terms are adjusted to be consistent with a particular choice of the cavity boundary. This adjustment is necessary because that boundary is intrinsically arbitrary, but the bulkelectrostatic contribution is very sensitive to it. The most accurate of the type 4 continuum solvation models are SM8 [129], SM8AD [128], and solvation model based on density (SMD) [127]. (These are sometimes called SMx models where x specifies which one.) The conductor-like screening model for real solvents (COSMO-RS) [132,133] adopts a different strategy in which a conductor-like screening calculation is performed on a molecule to generate a set of screening charges on the molecular cavity.

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The distribution of these charges forms a unique electrostatic fingerprint (called the σ-profile) that is characteristic of that molecule. The solvation free energy is then evaluated from a statistical mechanical procedure involving the interaction of the screening charges of the solute and those of the solvent. The COSMO-RS model has good accuracy (similar to the SMx models), at least for neutral solutes. The coupling of the solute to the solvent is directly related to Gibbs free energy change associated with the transfer of a particle in the gas phase to the solvent in a process in which the concentration in moles per liter does not change [134]. Therefore, it is sometimes convenient to use a standard state where the solute concentration in both phases is 1 mol L−1, and this standard state is denoted by “*” in ∆GS* , to distinguish it from ∆GS, which corresponds to a gas-phase partial pressure of the solute of 1 atm or 1 bar. We note that when metal complexes have open coordination sites, it is generally inaccurate to assume that a continuum solvation approach will accurately reflect the interactions of the metal with the “missing” first solvation shell. In principle, first-shell solvent molecules could be regarded as ligands that are explicitly included in the atomistic model. Indeed, for small, highly charged ions, it may be necessary for highest accuracy to include explicitly not only the first solvation shell but also the second [135,136]. However, inclusion of even the first shell raises questions about conformational averaging, and the best practical way to address these questions has not yet been convincingly demonstrated. The option of adding explicit solvent is more general than just filling open coordination sites. It has been concluded that continuum solvent models become quantitatively inaccurate near highly concentrated regions of charge [33,130]. Therefore, it was recommended that one should add a single explicit water molecule to any anion containing three or fewer atoms, to any anion with one or more oxygen atoms bearing a more negative partial atomic charge than the partial atomic charge on oxygen in water, and to any (substituted or unsubstituted) ammonium or oxonium ion [130]. Next, we comment on the issue of molecular geometry. Many solvation calculations use the gas-phase geometry in both phases. This is often reasonable because the difference in solvation energies calculated with gas-phase geometries and liquid-phase geometries is often less than other uncertainties in the calculations. However, it is safer to optimize the geometry separately in each phase. In cases where the conformational or other structural change associated with solvation is large, one can include this contribution to the solvation free energy computed on the solution-phase optimized geometry as follows: ∆GS ≅ ∆GS (soln geom) + Egas (soln geom) − Egas (gas geom)

(6.28)

Discussions of the use of gas-phase and solution-phase frequencies are given elsewhere [137,138]. 1. Absolute Potential of the Aqueous SHE In calculating free energies of solvation of ionic species (with charge ±z), a distinction is made between the absolute or intrinsic free energy of solvation and the real free energy of solvation, where the latter includes the contribution associated with the surface potential (χ) of the solvent [139]. The surface potential of water is controversial, and a rather large scatter of values, differing by more than 1 eV, has been reported [139–143]. The choice of χ directly affects the real solvation free energy of the proton and therefore also the value of ESHE, which is determined by the cycle in Figure 6.3: ESHE =

−∆Grxn F

  −∆Grxn = ∆Gion + ∆Gatom + ∆GS (H + ) = ∆ f G  (H + ) + ∆GS (H + )

(6.29a) (6.29b)

At present, the ESHE values of 4.28 and 4.42 V are most commonly used; these are derived from values of ∆GS* (H + ) of −1112.5 [144,145] and −1098.9 kJ mol−1 [140], respectively, in conjunction

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Organic Electrochemistry ΔGrxn H+(aq)

+

e–(g)

1/2H2(g)

–ΔGoatom

ΔG = 0

–ΔGso(H+)

-ΔGoion H+(g)

FIgURE 6.3

+

e–(g)

H (g)

Thermodynamic cycle for the SHE.

with a value of ΔfG°(H+) of 1517.0 kJ mol−1. The reader should note that the values of the two terms in Equation 6.29b depend on the choice of statistical formalism used to treat the electron, and the preceding values are based on Boltzmann statistics. The corresponding ∆GS* (H +) and ΔfG°(H+) values based on Fermi–Dirac statistics are −1108.9 [42], −1095.3, and 1513.3 kJ mol−1 [114]. The quantity ∆GS* (H +) is positively shifted by 3.6 kJ mol−1, and ΔfG°(H+) is negatively shifted by the same amount; therefore, the value of ESHE is independent of convention. The ESHE value of 4.42 V includes an estimate of the contribution due to the surface potential of water. More recent experimental estimates of ESHE (4.05, 4.11, and 4.21 V) [146–148] derived from nanocalorimetric measurements have been reported; however, the uncertainty associated with this technique is still relatively large. Because the total charge is conserved in a reaction, the contribution due to the surface potential cancels out in a chemically balanced chemical reaction that occurs in a single phase. As such, where calculation of equilibrium reduction potentials involving a single phase is concerned, it should not matter whether the contribution from surface potential is included in the solvation free energy, as long as this is done consistently for all reacting species and products. This raises the question as to whether continuum solvent models are designed to predict real or absolute solvation free energies. Continuum solvent models generally contain parameters (e.g., atomic radii used to construct the molecular cavity) that have been optimized to reproduce experimental solvation free energies. However, the experimental solvation free energies of ionic solutes are indirectly obtained via thermochemical cycles involving, for example, the solvation free energy of the proton, aqueous pKa values, and gas-phase reaction energies. Accordingly, the ESHE values that should be used with a continuum model are those that are based on a consistent ΔG S(H+). Table 6.2 provides an overview of several continuum solvent models typically used in aqueous calculations, the ΔGS(H + ) upon which they are based, and examples of the levels of theory for which they have been most extensively benchmarked. As shown, some continuum solvent models such as the (C) PCM-UAHF and (C)PCM-UAKS models, where UAHF denotes the use of united-atom parameters optimized for Hartree-Fock calculations, and UAKS denotes the use of united-atom parameters optimized for Kohn-Sham calculations, are based on ∆GS* (H + ) values that are slightly different from those used to derive the ESHE values of 4.28 and 4.42 V. In such cases, where the difference is significant, one could adjust the value of the ESHE to make it compatible with the continuum solvent model as shown in Table 6.2. The COSMO-RS model was parameterized using solvation free energies (and related data) of neutral solutes [133], and therefore its compatibility with a particular ESHE is unclear. 2. Nonaqueous Systems In nonaqueous solution, there is no primary reference electrode equivalent to the aqueous SHE or SCE. Nonaqueous silver electrodes using silver nitrate or perchlorate are reliable reference electrodes for nonaqueous solutions; however, details on the actual Ag+ concentration or salt anion in the Ag+/Ag are often not reported, making it difficult to directly compare potentials obtained from different studies [54]. Although aqueous reference electrodes are often used for nonaqueous systems, the liquid junction potential between the aqueous and nonaqueous solutions can affect the measurements. For these reasons, the IUPAC Commission on Electrochemistry has recommended

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TAbLE 6.2 Examples of Commonly Used Solvent Models and the Levels of Theory at Which They Are Applied ∆GS* (H+ ) (kJ mol−1)

Level of Theory

ESHE (V)

(C)-PCM-UAHF [149] (C)-PCM-UAKS SM6 [130]

−1093.7 −1093.7a −1105.8

4.47 4.47 4.34

SMD [127]

−1112.5

SM8 [129] and SM8AD [128]

−1112.5

HF/6-31G(d) for neutrals and HF/6-31+G(d) for ions B3LYP or PBE0/6-31+G(d) MPW25/MIDI!6D or 6-31G(d) or 6-31+G(d) B3LYP/6-31+G(d,p) B3PW91/6-31+G(d,p) and any DFT method that can deliver a reasonably accurate electronic density for the solute of interest Any electronic structure model delivering a reasonable continuous density distribution HF theory and many local and hybrid density functionals with basis sets of up to minimally augmented polarized valence double-zeta quality BP/TZP

Solvent Model

COSMO-RS [133]

—

4.28 4.28

—

The value of the solvation free energy of the proton upon which the model is based and corresponding aqueous ESHE values are also shown. a Assumed value.

that the ferrocenium/ferrocene (Fc+/Fc) couple be used as an internal reference for reporting electrode potentials in nonaqueous solutions [150], and knowledge of its absolute potential is therefore essential for calculations to be referenced to this electrode. The absolute potential of the Fc+/Fc couple in a nonaqueous solvent can be quite simply obtained from ESHE and the conversion constant between aqueous SHE and (Fc+/Fc) in a nonaqueous solvent. Pavlishchuk and Addison determined the conversion constants between various reference electrodes, including the Fc+/Fc couple in acetonitrile and aqueous SCE (and SHE) [54]. Thus, using ESHE values of 4.28 and 4.42 V in conjunction with the conversion constant of 0.624 V leads to Fc+/Fc potentials of 4.90 and 5.04 V, respectively. More recent calculations using the SMD and COSMO-RS solvent models (in conjunction with gas-phase free energies calculated at G3(MP2)-RAD-Full-TZ and Fermi–Dirac statistics for the electron) provided estimates of 4.96 and 4.99 V for the Fc+/Fc potential in acetonitrile, respectively [151]. These values are generally in good agreement with the two “experimental” values of the Fc+/Fc potential (within a 100 mV). The choice of Fc+/Fc potential for continuum-solvent-based predictions is less obvious, and one could instead adopt an approach analogous to cycle B in Figure 6.2 where both half-cells are treated using the same continuum solvent model. Related to this point, the reader should note that not all solvent models have been designed to predict solvation free energies in nonaqueous solvents. Examples of models that have been designed to treat nonaqueous solutions are the SMD [127] and the COSMO-RS models [132,133]. The PCMUAKS and PCM-UAHF models were designed specifically for predicting aqueous free energies of solvation [149], although there have been attempts [152] to extend these models to nonaqueous solvents through the manipulation of other parameters within the solvent model such as the scaling factor (α) that relates to the solvent-inaccessible cavity.

C.

STANDARD STATES

When calculating solution-phase reaction energies using a thermodynamic cycle that combines quantities obtained from different sources and/or calculations, it is important to pay attention to the standard state of these quantities. The literature on calculating solvation free energies by quantum

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mechanics usually uses a solute standard-state concentration of 1 mol L−1, whereas 1 molal is more common in some other subfields of chemical thermodynamics. The approximation of molality by molarity is reasonable for aqueous solutions since the density of water is approximately 1 kg L−1 for quite a large range of temperatures. This is not necessarily true for solutions involving organic solvents since the density of these solvents is typically much lower. As noted earlier, the quantity yielded directly by continuum solvation models without a concentration term is the Gibbs free energy change associated with the transfer of a particle in the gas phase to the solvent, where the molarity of the solute is the same in both phases. On the other hand, gas-phase thermodynamic quantities are conventionally calculated using a standard state of 1 atm. The conversion between free energies of solvation in the two conventions is straightforward when we recall the standard states are actually ideal gases and ideal solutions. Thus, the standard-state quantities correspond to measurements at infinite dilution followed by extrapolation to unit activity as if the activity coefficient were unity (ideal behavior). Therefore,  ∆GS = ∆GS* + ∆Gconc

(6.30)

where  RT   ∆Gconc = RT ln    P 

(6.30a)

where R is the gas constant P° is the standard-state pressure   At 298 K, we get ∆Gconc = 7.96 kJ/mol for P° = 1 bar and ∆Gconc = 7.93 kJ/mol for P° = 1 atm. A separate issue relating to standard states is that experimental measurements are not usually made at either an activity of one or a molarity of one. For example, they may be made in systems buffered to keep particular reactant and/or product concentrations at some convenient concentration. For example, reductive chlorination potentials are nearly always measured with the chloride ion concentration at about 10−3 M—these are conditional potentials, but they are not standard or formal potentials; however, they can be converted to standard concentrations. Similarly, to use thermodynamic data in applications, one must convert from tabulated standard-state quantities to quantities pertaining to real experimental conditions. To facilitate the comparison between standard free energies and those pertaining to nonstandard conditions, we note that the Gibbs free energies of reaction at nonstandard concentrations and those at standard concentrations are related by

Q  ∆G = ∆G  + RT ln    Q 

(6.31)

where Q is the reaction quotient ΔG and Q are for nonstandard concentrations ΔG° and Q° are for standard states At equilibrium, ΔG = 0 and Q becomes the equilibrium constant K, so this yields  K  ∆G  = − RT ln    Q 

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(6.31a)

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D. RATES OF ELECTRON TRANSFER The focus of this chapter is on the prediction of standard reduction potentials, and not on kinetics, but we note here that the sum of two standard half reactions defines the standard driving force ΔG° for an electron-transfer reaction between a donor D and an acceptor A. For convenience of notation, we will here write D and A as neutral species and the post-electron-transfer products D + and A− as singly positively and negatively charged species, respectively, but there is no restriction on the initial and final charge states beyond the obvious one that after a single electron-transfer D will be one unit more positively charged and A one unit more negatively charged. In the Marcus theory, the driving force is a key variable for the prediction of free energies of activation associated with electron-transfer reactions. This free energy of activation can be used in a transition-state theory equation or a diabatic collision theory approach to compute rate constants for electron-transfer reactions. In particular, the Marcus theory [153] takes the free energy of activation to be

( λ + ∆G ) = 

∆G

‡

2

4λ

(6.32)

where we have omitted some work terms necessary to bring the reagents together and where λ is the reorganization energy associated with the electron-transfer reaction. The reorganization energy may be taken as the sum of two components: an outer-sphere and an inner-sphere reorganization energy. The former is associated with the change in solvation free energy that occurs when a generalized bulk solvent coordinate equilibrated with the pre-electron-transfer state is confronted instantaneously with the post-electron-transfer state. Such changes in solvation free energy may be computed using two-time-scale continuum solvation models [154–156] that permit the fast (optical) component of the solvent reaction field to be equilibrated to the post-electron-transfer state while the slow (bulk) component remains frozen as it was equilibrated to the pre-electron-transfer state. The free energy of solvation of the charge-transfer (CT) state interacting with the nonequilibrium two-time-scale reaction field minus the free energy of solvation of the pre-CT state interacting with its fully equilibrated reaction field defines the outer-sphere reorganization energy. The inner-sphere reorganization energy, on the other hand, is associated with changes in the donor and acceptor structures (including possibly their first solvation shells) as they relax following the electron transfer. From a computational standpoint, these various quantities are readily computed. Thus, for instance, by computing the energy change as D+ relaxes from the geometry of D to that of D+ (which in some instances may involve including the first solvation shell of D/D+), one may compute the contribution of the donor molecule to the inner-sphere reorganization energy. Since kinetics is a digression from our main subject, we will not develop this topic further, but we emphasize that the computational techniques outlined here to compute electron-transfer driving forces, combined with approaches to compute reorganization energies, offer a practical avenue to addressing electrontransfer rate questions.

IV. EXAMPLES This section contains examples of calculations of reduction potential. All calculations were performed using Gaussian 09 [157] or MOLPRO 2009 [158].

A.

AQUEOUS STANDARD ONE-ELECTRON REDUCTION POTENTIALS NITROXIDES AND QUINONES

OF

In this example, we calculate the standard potentials of the aqueous one-electron reduction half reactions shown in Figure 6.4.

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COOH

COOH 4-COOH-TEMPO

+

e–

0.81 N

– –

–

+ N O

O CONH2

CONH2 3-CONH2-TCPO

0.96 –

N

O

O

O

O +

benzoquinone

0.10 O

O

O +

O

–

e–

O

2,3-dimethylnaphthoquinone

FIgURE 6.4

+

e–

– –

+ N

expt E o (V) vs SHE

–

e–

–0.24 O

Species studied with their experimental reduction potentials (see Table 6.3 for details).

The relevant computational data are shown in Table 6.3. The gas-phase Gibbs free energies were computed at the G3(MP2)-RAD(+) level of theory, which is a modification of the G3(MP2)-RAD [71] method. The (+) signifies that calculations originally defined to involve the 6-31G(d) basis set have been carried with the 6-31+G(d) basis set so as to allow for an improved description of anionic species.  , is calculated using cycle A in Figure 6.2: The aqueous-phase Gibbs free energy of reaction, ∆Gsoln     ∆Gsoln = Ggas (Red) − Ggas (Ox) − Ggas (e) + ∆GS (Red) − ∆GS (Ox)

(6.33)

 By substituting the appropriate values into this expression, one obtains the ∆Gsoln in Table 6.3 and the corresponding standard reduction potentials. The values of 4.47 and 4.28 V for ESHE were used in conjunction with calculations employing the CPCM-UAHF and SMD solvent models as outlined in Table 6.2. The table shows that while the approach performs reasonably well for nitroxides, its performance is much less satisfactory for the quinones where the magnitude of the errors is 380 mV or larger for both solvent models. This example illustrates the difficulty associated with the direct calculation of absolute reduction potentials where performance depends heavily on the accuracies of absolute solvation free energies of the reactants and products. In particular, all half reactions generate or consume a charged species, and because the uncertainty in the solvation free energies associated of these species are significantly higher, this directly impacts the accuracy of absolute potentials. The present example also illustrates that the good performance of directly calculated reduction potentials by a given method for a particular class of compounds does not necessarily extend to other types of compounds. An interesting observation for the four cases in Table 6.3 is that in every instance, the reduced product would be expected to be a much stronger hydrogen bond acceptor than the oxidized precursor. Thus, first-solvent shell water molecules are very important. An alternative approach is to calculate relative reduction potentials, which can be more accurate by systematic error cancellation. For example, the data in Table 6.3 reveal that calculations based on the CPCM-UAHF model underestimate the standard potentials for quinones by about 600 mV. Such a systematic error will largely cancel out for the reaction shown in Figure 6.5. The potential associated with this reaction is readily obtained from the data in Table 6.2 as the reduction potential of 2,3-dimethylnaphthoquinone less that of benzoquinone. Using the

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4-COOH-TEMPOg

H0 (kJ mol−1) ΔGtherm (kJ mol−1) G°gas (kJ mol−1)b ∆G°S (UAHF; kJ mol−1)c ∆G°S (SMD; kJ mol−1)c ∆G°soln (UAHF; kJ mol−1)d ∆G°soln (SMD; kJ mol−1)d E° (UAHF) rel. SHEe (V) E° (SMD) rel. SHEf (V) E° (expt)

3-CONH2-TCPOg

benzoquinone

2,3-Dimethylnapthoquinone

Ox

Red

Ox

Red

Ox

Red

Ox

Red

−1,761,669.5 −108.5 −1,761,778.0 −233.2 −242.6

−1,762,348.8 −110.2 −1,762,459.0 −42.2 −44.7 −486.3

−1,603,319.8 −105.6 −1,603,425.4 −212.3 −241.9

−1,604,000.7 −107.6 −1,604,108.3 −39.9 −52.5 −506.9

−1,000,069.2 −78.5 −1,000,147.8 −28.1 −24.9

−1,000,250.1 −79.1 −1,000,329.2 −243.3 −233.0 −393.1

−1,608,881.0 −102.5 −1,608,983.5 −13.5 −19.7

−1,609,042.0 −101.0 −1,609,143.0 −209.9 −215.6 −352.2

−479.5 0.57 (−0.24) 0.69 (−0.12) 0.81 [159]

−489.9 0.78 (−0.18) 0.80 (−0.16) 0.96 [159]

−386.0 −0.40 (−0.50) −0.28 (−0.38) 0.10 [160]

−351.8 −0.82 (−0.58) −0.63 (−0.39) −0.24 [160]

Theoretical Calculation of Reduction Potentials

TAbLE 6.3 Computational Data for the Calculation of Standard Reduction Potentials at 298 K and Relative to SHEa

Signed errors are shown in parentheses. The gas-phase energies were computed at the G3(MP2)-RAD(+) level. Solvation calculations using the CPCM-UAHF and SMD models were performed by the HF/6-31+G(d) and B3LYP/6-31+G(d) methods on the respective solution-phase optimized geometries. CPCM-UAHF solvation free energies were performed at the ROHF/6-31+G(d) level on UHF/6-31+G(d) solution optimized geometries for open-shell species. b G o = H + ∆G gas 0 therm . c Solvation free energies printed in Gaussian 09 correspond to ∆G * and Equation 6.30 is used to obtain ∆G °. S S d ∆G°soln calculated from Equation 6.32. e E SHE = 4.47 V. f E SHE = 4.28 V. g 4-COOH-TEMPO = 2,2,6,6-tetramethylpiperidinoxyl; 3-CONH -TCPO = 2,2,5,5-tetramethyl-3-carbamido-3-pyrroline-1-oxyl. 2 a

249

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250

Organic Electrochemistry O

O

–

O

–

O

+

+

O

FIgURE 6.5

O

O

O

An isodesmic CT reaction. –

A (soln) + Ref (soln)

ΔGCT

–

– –ΔGso(A) –ΔGso (Ref )

A(g)

ΔGso(Ref )

ΔGso (A ) –

+

–

A (soln) + Ref (soln)

Ref (g)

ΔGogas

–

A (g)

+ Ref (g)

FIgURE 6.6 Thermodynamic cycle for a CT reaction.

CPCM-UAHF model, this CT potential is −0.42 V. Thus, by using benzoquinone as a reference molecule for which the experimental standard potential is known (0.10 V), one can estimate the standard potential of 2,3-dimethylnaphthoquinone by adding the CT potential to E°(benzoquinone) to give E°(2,3-dimethylnaphthoquinone) = −0.32 V. This approach brings the error down from 580 to 80 mV. More generally, for the CT reaction between A and a reference molecule (Ref) with known E°, the standard potential E°(A/A−) may be obtained from the thermodynamic cycle in Figure 6.6 and Equation 6.34a:  ∆GCT = ∆Ggas + ∆GS (A• − ) + ∆GS (Ref ) − ∆GS (A) − ∆GS (Ref • − )

(

)

E  A A• − =

∆GCT  + Eexpt Ref Ref • − 96.5 C mol −1

(

)

(6.34) (6.34a)

 An added advantage of this approach is that ESHE is no longer needed, thereby eliminating a source of uncertainty. However, since the method relies on systematic error cancellation, it is expected to work best when the reference molecule is structurally similar to A. The major limitation of this approach is that a structurally similar reference with accurately known E° may not always be available.

B.

CHEMICALLY IRREVERSIBLE PROCESSES—REDUCTIVE DECHLORINATION

Next, we show how the reduction potentials corresponding to the dissociative electron-transfer reactions of some alkyl halides in aqueous and nonaqueous solutions (Figure 6.7) are calculated. The relevant computational data and results are presented in Tables 6.4 and 6.5, respectively. Since the potentials of reactions 1, 3, and 4 are measured in DMF and are referenced to the aqueous SCE, a 0.172 V [53] correction for a liquid junction potential was applied to the calculations. Accordingly, using the reductive cleavage of carbon tetrachloride (reaction 3) as example, its reduction potential was calculated as follows: ∆Gsoln = − 361.4 kJ mol −1 E =

−∆Gsoln  − ESHE − E  ( SCE SHE ) − E j 96.5

= 3.75 − 4.28 − 0.241 + 0.172 = −0.60 V

(6.35)

where the calculations are referenced to the aqueous SCE and E°(SCE/SHE) is its potential relative to aqueous SHE (0.241 V) [42].

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251

Theoretical Calculation of Reduction Potentials E (expt) –

Cl(soln) + e

Cl (soln)

CCl4(soln) + e– CHCl3(soln) + e– Cl

Cl Cl (soln) + e–

Cl

Cl Cl Cl Cl Cl Cl

CCl3(soln) + Cl–(soln)

E°DMF = –0.58 vs SCEaq

CHCl2(soln) + Cl–(soln)

E°DMF = –0.84 vs SCEaq

Cl Cl

Cl

Cl

E°aq = 2.59 vs SHE; E°DMF = 2.12 vs SCEaq

–

Cl Cl (soln) + Cl–(soln)

Eaq = 0.11 vs SHE; [Cl–] = 10–3 M

(soln) + 2Cl–(soln)

Eaq = 1.15 vs SHE; [Cl–] = 10–3 M

Cl

Cl – Cl (soln) + 2e Cl

Cl

Cl

Cl

Cl

Cl

Cl Cl (soln) + H+(soln) + 2e–

Cl

Cl

Cl

Cl Cl (soln) + Cl–(soln) H

Eaq = 0.67 vs SHE; [Cl–] = 10–3 M, [H+] = 10–7 M

FIgURE 6.7 Species studied with their experimental reduction potentials in V (see Table 6.5 for details).

TAbLE 6.4 Calculated gas-Phase gibbs Free Energies and Solvation Free Energies at 298 Ka Cl•

Cl−/Cl−⋅H2O

CCl4

CCl•3

CHCl•2

CHCl3

G°gas (kJ mol−1)a −1,206,953.9c −1,207,306.1/−1,407,825.0 −4,928,225.9 −3,721,026.3 −3,722,725.7 −2,515,500.8 ∆GS° (SMD) H2Ob 8.4 —/−260.3 — — — — 2.1 −12.9 −0.8 1.8 −264.8/— −4.2 ∆GS° (SMD) DMFb C2Cl5•

C2Cl6 ° (kJ mol ) −7,442,299.2 Ggas ∆GS° (SMD) H2Ob 13.2

−6,235,086.3 13.9

−1 a

a b c

C2HCl5

C2Cl4

−6,236,794.9 −5,028,088.6 3.0 15.7

H+

H2O

−26.3 −1,104.6

−200,483.0 —

Computed at the G3(MP2)-RAD(+) level of theory. Calculations (in kJ mol−1) performed by the B3LYP/6-31+G(d) method on solution-phase optimized geometries. Includes spin–orbit correction (−1.34 millihartrees).

As mentioned in Section III.B, first-solvent shell interactions are likely to be very important for species with regions of concentrated charge such that a continuum model is likely to be inadequate. The reader should therefore note that the SMD, SM6, and SM8 solvent models are to be used as mixed discretecontinuum models in such cases; in particular, they have been parameterized to reproduce the experimental aqueous solvation free energy of the Cl−⋅H2O cluster and (H 2O)2 dimer, not the solvation free energy of bare Cl− or H2O [33,127,129,130]. As such, for the aqueous reactions that involve a bare chloride ion, that is, reactions 2 and 5 to 7, the calculations were carried out with the addition of a water of hydration,  as shown in Table 6.5. Using the last reaction as example, the calculated ∆Gsoln was obtained as follows:   = ∆Ggas + ∆∆GS = − 961.5 kJ mol −1 ∆Gsoln

(6.36)

where ∆∆GS = ∆GS (Cl ⋅ H 2O − ) + ∆GS (C2HCl 5 ) − ∆GS (H + ) − ∆GS (C2Cl6 ) − ∆GS (H 2O) = 842.6 kJ mol −1

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(6.37)

252

Organic Electrochemistry

TAbLE 6.5 Calculated Reduction Potentials and Experimental Valuesa E/ V (calc)

E/ V (expt)

1 2 3

Cl•(dmf) + e− → Cl−(dmf) Cl•(aq) + H2O(l) + e− → [Cl(H2O)]−(aq) CCl 4 (dmf ) + e − → CCl3• (dmf ) + Cl − (dmf )

2.03 2.40c −0.60

2.12 [161] 2.59 [161] −0.58 [162]

4

CHCl3 (dmf ) + e − → CHCl•2 (dmf ) + Cl − (dmf )

−0.93

−0.84 [162]

5 6 7

C2 Cl6 (aq) + H 2 O(l) + e − (g) → C2 Cl•5 (aq) + [Cl(H 2 O)]− (aq) C2Cl6(aq) + 2H2O(l) + 2e−(g)→C2Cl4(aq) + 2[Cl(H2O)]−(aq) C2Cl6(aq) + H2O(l) +  H + (aq) + 2e−(g)→C2HCl5(aq) + [Cl(H2O)]−(aq)

−0.20b,c  0.91b,c 0.58b,c

a b c

0.11b [1] 1.15b [1] 0.67b [1]

Reactions in DMF and aqueous solution are referenced to SCE(aq) and SHE(aq), respectively. These potentials correspond to the experimental conditions [Cl−] = 10−3 mol L−1 and pH = 7. Calculations that include an explicit water of hydration. The experimental solvation free energy of the water (−8.6 kJ mol−1) that corresponds to a standard state of [H2O] = 55 mol L−1 (i.e., pure water) and 1 atm in the liquid and gas phase was used in these calculations.

Note that in these calculations, we have used the experimental value for the solvation free energy for water (−8.6 kJ mol−1) [163] under the conventional standard state for pure liquids, that is, mole fraction of 1 in the liquid phase and 1 atm in the gas phase. In these reactions, the experimental potentials for the reductive cleavage of hexachloroethane were referenced to SHE, and therefore no correction for Ej was applied. However, the potentials corresponded to nonstandard conditions of [Cl−] = 10−3 mol L−1 and pH 7, and a correction using Equation 6.31 was applied to arrive at the values in Table 6.5:

(

 [1 M C2HCl 5 ][10 −3 M Cl − ] / [1 M C2Cl6 ][10 −7 M H + ]  ∆Gsoln = ∆Gsoln + RT ln   [1 M C2HCl 5 ][1 M Cl − ] / [1 M C2Cl6 ][1 M H + ] 

(

)

 = ∆Gsoln + RT ln(10 4 ) = −938.7 kJ mol −1

)    (6.38)

Accordingly, the potential for this two-electron reduction is

E=−

C.

−938.7 − 4.28 = 0.58 V 2 × 96.5

(6.39)

CONSTRUCTING A POURBAIX DIAGRAM FOR THE TWO-ELECTRON REDUCTION OF O-CHLORANIL

Consider the two-electron reduction of o-chloranil (OCA) in aqueous solution [164]. Depending on the pH of the solution, the reduction process can be represented in one of the following ways as shown in Figure 6.8. The corresponding standard reduction potentials are denoted E°(OCA/OCA2−), E°(OCA,H+/ OCAH−), and E°(OCA,2H+/OCAH2), and these are related to each other as follows:

(

)

(

)

E  OCA OCA 2 − = E  OCA, H + OCAH − +

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RT ln K 2 2F

(6.40)

253

Theoretical Calculation of Reduction Potentials Cl

Cl

Cl

O

O –

Cl

O –

2e–

+ Cl

Cl

O Cl

Cl

Cl

Cl O

Cl

+

2e–

OH

Cl

O –

+ H+

O

Cl

Cl

Cl

Cl

Cl

Cl

Cl

O + 2e– + 2H O

Cl

OH

Cl

+

OH

Cl

Cl

FIgURE 6.8

Cl

The microspecies present in the two-electron reduction of OCA in aqueous solution.

(

)

(

)

E  OCA OCA 2 − = E  OCA, 2H + OCAH 2 +

RT ln K1K 2 2F

(6.41)

where K1 and K2 are the first and second acid dissociation constants of OCAH2. From Equation 6.11, the potential for the E(OCA,2H+/OCAH2) is

(

)

E = E °′ OCA, 2H + OCAH 2 +

RT [OCA 2 − ][H + ]2 ln 2F [OCAH 2 ]

(6.42)

Equation 6.42 can alternatively be expressed in terms of the acid dissociation constants (K1 and K2) of the conjugate acid of the reduced product (H2A):

(

)

E = E °′ OCA, 2H + OCA 2 − +

RT RT S ln K1K 2 + K1[H + ] + [H + ]2 + ln Ox 2F 2 F SRed

(

)

(6.43)

SOx = [OCA]

(6.43a)

SRed = [OCA 2− ] + [OCAH − ] + [OCAH 2 ]

(6.43b)

Using techniques such as cyclic voltammetry, one can measure a half-wave potential (E1/2) where the concentrations of the reductant are approximately equal to the oxidant, that is, SOx = SRed, and Equation 6.43 becomes

(

)

E1/ 2 = E °′ OCA, 2H + OCA 2 − +

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RT ln K1K 2 + K1[H + ] + [H + ]2 2F

(

)

(6.44)

254

Organic Electrochemistry

TAbLE 6.6 Calculateda Reduction Potentials in V and pKa Values 0.83 (0.79) [164] 0.63 (0.67) [164] 0.41b (5) [164] 9.2c

E°(OCA,2H + /OCAH2) E°(OCA,H + /OCAH−) E°(OCA/OCA2−) pK1 pK2

Experimental values, where available, are shown in parentheses. a Calculations are based on the G3(MP2)-RAD(+) gas-phase energies with SMD solvation energies obtained at the B3LYP/6-31+G(d) level and ESHE of 4.28 V. b Calculated from Equation 6.40 using the data in this table. c Calculated using a proton-exchange method [34,35] using orthoquinone (expt pK = 13.4) a [165] as the reference.

From the calculated reduction potentials in Equations 6.40 and 6.41 as well as the acid dissociation constants (K1 and K2) of the diprotic acid, OCAH2, a chemical speciation plot denoting the dominant microspecies in a particular pH range can be obtained. The data needed for such a plot are shown in Table 6.6. From Equation 6.43, three distinct linear pH ranges can readily be identified. In the range where pH < pK1, [H+] ≫ K1 ≫ K2, the molecule OCAH2 is the predominant form of the reduced product, and the midpoint potential has a pH dependence based on Equation 6.44:

(

)

E1/ 2 = E °′ OCA, 2H + OCAH 2 +

RT ln[H + ]2 2F

(6.45)

In the other two linear segments at pK1 < pH < pK2 and pH > pK2, the reduced product exists predominantly as OCAH− and OCA2−, respectively, and the corresponding half-wave potentials have pH dependence following equations:

(

)

E1/ 2 = E °′ OCA, 2H + OCAH 2 +

(

)

RT ln( K1[H + ]) 2F

(6.46)

RT ln( K1K 2 ) 2F

(6.47)

E1/ 2 = E °′ OCA, 2H + OCAH 2 +

Extrapolation of the three linear segments (with theoretical slopes −2.303mRT/2F, where m is the number of protons involved in the reaction) to pH 0 yields the formal potential E°′(OCA,2H + /OCAH2), E°′(OCA,H+/OCAH−), and E°′(OCA/OCA2−), respectively. Collectively, this information can be used to construct a E versus pH (Pourbaix diagram) as shown in Figure 6.9. The vertical lines correspond to the pKas of the diprotic OCAH2 acid. The reader should note that the formal potential E°′ is pH invariant since the condition [H+] = 1 mol L −1 applies. However, half-wave potentials are strongly pH dependent, and these are quite often reported instead of standard or formal reduction potentials. Thus, in comparing with experiment, it is also important to examine the details of the experimental measurement to ascertain whether the calculation corresponds to the same quantity as the one reported.

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255

Theoretical Calculation of Reduction Potentials pH < pK1

pK1 < pH < pK2

pH > pK2

0.9 E°(OCA, 2H+/OCAH2)

Half-wave potential (V) versus SHE

0.8 0.7

E°(OCA, H+/OCAH–)

0.6 0.5

E°(OCA/OCA2–)

0.4 0.3

OCA/OCAH2

0.2

OCA/OCAH–

OCA/OCA2–

0.1 0 0

2

4

6

8

10

12

14

pH

FIgURE 6.9 An E versus pH diagram (Pourbaix diagram) for OCA. The vertical dotted lines correspond to the pKas of OCAH2 and indicate the pH regions in which various stable species predominate.

V. CONCLUDINg REMARKS We have presented an introductory guide to carrying out QM continuum solvent prediction of solution-phase reduction potentials. We stress that reduction potentials are equilibrium thermochemical parameters. We discussed issues pertaining to thermochemical conventions for the electron, the choice of standard electrode, and the advantages and limitations of methods based on thermodynamic cycles for calculating reduction potentials. Just as in experimental work, a key consideration for predicting chemically accurate reduction potentials is the difficulty of obtaining accurate estimates of the solvation free energies of ionic species. Careful work often involves including (or expanding) a first solvation shell, particularly in solvents donating or accepting strong hydrogen bonds. Relative reduction potential calculations can partly remedy this problem by exploiting systematic error cancellation in the solvation calculations.

ACKNOWLEDgMENTS We gratefully acknowledge support from the Australian Research Council under their Centres of Excellence program and the generous allocation of computing time on the National Facility of the National Computational Infrastructure. MLC also acknowledges an ARC Future Fellowship. This work was supported in part by the U.S. Army Research Lab under Grant No. W911NF09-1-0377 and the U.S. National Science Foundation under Grant No. CHE09-56776.

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Section II General Preparative Aspects

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7

Preparative Electrolysis on the Laboratory Scale Jakob Jörissen and Bernd Speiser

CONTENTS I.

II.

Principles of Electroorganic Cell Operation ........................................................................ 265 A. Introduction ................................................................................................................... 265 B. Basic Requirements and Theoretical Definitions .........................................................266 1. Precondition of Any Electrolysis: A Closed Electric Circuit ................................266 2. Types of Electrochemical Reactions ...................................................................... 266 3. Cell Current ........................................................................................................... 266 4. Electrode Potential ................................................................................................. 267 5. Cell Voltage............................................................................................................ 270 6. Basic Definitions for Chemical Reactions, Also Valid for Electroorganic Electrolysis .....271 7. Operation Modes of (Electrochemical) Reactors .................................................. 271 8. Control of Electrochemical Reactions ................................................................... 272 Components of Electroorganic Reaction Systems ................................................................ 275 A. Electrodes...................................................................................................................... 275 1. General Requirements for Electrode Materials ..................................................... 275 2. Special Requirements for Cathode Materials ........................................................ 276 3. Special Requirements for Anode Materials ........................................................... 276 B. Examples of Electrode Materials .................................................................................. 277 1. Platinum, Platinum Metals, Other Noble Metals, and Their Alloys ..................... 277 2. Nickel ..................................................................................................................... 277 3. Iron (Mild Steel), Stainless Steel ........................................................................... 278 4. Lead ....................................................................................................................... 278 5. Mercury.................................................................................................................. 278 6. Carbon.................................................................................................................... 279 7. Conductive Ceramic, for Example, Ebonex ........................................................... 279 8. Electrocatalytic Coatings on Carrier Materials .....................................................280 9. Boron-Doped Diamond Coating ............................................................................280 C. Examples of Electrode Types and Their Special Properties ........................................ 281 1. Porous vs. Smooth Electrodes................................................................................ 281 2. Gas-Evolving Electrodes........................................................................................ 281 3. Sacrificial Anodes .................................................................................................. 281 4. Gas Diffusion Electrodes ....................................................................................... 282 D. Electrolytes ................................................................................................................... 282 1. Solvents .................................................................................................................. 282 2. Supporting Electrolytes ......................................................................................... 282 3. Solid Polymer Electrolyte Technology .................................................................. 283 E. Cell Separators .............................................................................................................. 283 1. Porous Materials .................................................................................................... 283 2. Ion Exchange Membranes ......................................................................................284 263

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F.

Electrochemical Cells ................................................................................................... 285 1. Homogeneous Current Density .............................................................................. 285 2. Electrode Potential Measurement ..........................................................................287 3. Uniform Mixing and Mass Transport .................................................................... 289 4. Temperature Control .............................................................................................. 289 5. Cell Construction Materials ................................................................................... 290 6. Procedure of Electrolysis Experiments, Mass Balancing and Charge Balancing...... 292 G. Examples of Electrochemical Cells .............................................................................. 294 1. H-Cell..................................................................................................................... 294 2. Beaker Glass Cells .................................................................................................294 3. Flow-Through Cells ............................................................................................... 296 4. Innovative Cell Constructions................................................................................ 298 References ...................................................................................................................................... 298 III Appendices ...........................................................................................................................300 III.A Appendix A: Solvents for Electrolysis .................................................................................300 III.A.1 General Considerations .........................................................................................300 III.A.1.a Proton Activity .....................................................................................300 III.A.1.b Usable Potential Range ........................................................................303 III.A.1.c Dielectric Constant...............................................................................303 III.A.1.d Dissolving Power .................................................................................303 III.A.1.e Temperature Range and Other Factors.................................................305 III.A.1.f Purification...........................................................................................306 III.A.2 Protic Solvents .......................................................................................................306 III.A.2.a Acid Solvents ....................................................................................... 306 III.A.2.b Neutral Solvents ...................................................................................308 III.A.2.c l,l,l,3,3,3-Hexafluoropropan-2-ol .........................................................309 III.A.2.d Basic Solvents ......................................................................................309 III.A.3 Aprotic Solvents .................................................................................................... 311 III.A.3.a Acetonitrile ..........................................................................................312 III.A.3.b Dimethylformamide .............................................................................313 III.A.3.c N-Methylpyrrolidone ...........................................................................314 III.A.3.d 3-Methyl-2-Oxazolidinone (3M2O) ..................................................... 314 III.A.3.e Hexamethylphosphoramide (HMPA) .................................................. 314 III.A.3.f Pyridine ................................................................................................315 III.A.3.g Dimethyl Sulfoxide (DMSO) ...............................................................315 III.A.3.h Sulfolane ..............................................................................................316 III.A.3.i Propylene Glycol Sulfite ......................................................................316 III.A.3.j Nitromethane........................................................................................316 III.A.3.k Nitrobenzene ........................................................................................317 III.A.3.l Propylene Carbonate (PC) ...................................................................317 III.A.3.m Benzene and Chlorobenzene ................................................................317 III.A.3.n Ethers ...................................................................................................317 III.A.3.o Methylene Chloride and 1,1,2,2-Tetrachloroethane.............................318 III.A.3.p Sulfur Dioxide...................................................................................... 318 III.A.4 Salts ....................................................................................................................... 318 III.A.5 Supercritical Fluids ............................................................................................... 319 References for Appendix A ............................................................................................................ 320 III.B Appendix B: Electrolytes...................................................................................................... 325 III.B.1 Anions ................................................................................................................... 325 III.B.1.a Perchlorate ...........................................................................................325 III.B.1.b Tetrafluoroborate .................................................................................. 325

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III.B.1.c Hexafluorophosphate, Hexafluoroarsenate ..........................................326 III.B.1.d Trifluoromethanesulfonate ...................................................................326 III.B.1.e Nitrate ..................................................................................................326 III.B.1.f Aromatic Sulfonates.............................................................................326 III.B.1.g Carboxylate Ion....................................................................................326 III.B.1.h Tetramethylaluminate and Tetraphenylborate ..................................... 327 III.B.2 Cations ................................................................................................................... 327 III.B.2.a Lithium Ions .........................................................................................327 III.B.2.b Sodium Ions .........................................................................................327 III.B.2.c Magnesium Ions ...................................................................................327 III.B.2.d Tetraalkylammonium Ions ...................................................................327 III.B.2.e Sulfonium Salts ....................................................................................328 III.B.2.f Cryptates ..............................................................................................328 III.B.2.g Polyelectrolytes .................................................................................... 328 III.B.3 Buffers ................................................................................................................... 328 III.B.3.a Acids ....................................................................................................328 III.B.3.b Bases ....................................................................................................329 III.B.3.c Buffer Systems ..................................................................................... 329 References for Appendix B ............................................................................................................ 329

I. A.

PRINCIPLES OF ELECTROORgANIC CELL OPERATION INTRODUCTION

Any electrolysis process needs a careful selection of optimal operation conditions as well as of an appropriate electrochemical cell. These decisions will be responsible for the result. The variety of possibilities is very broad and precise planning and usually also experiments are indispensable. The theoretical background of electroorganic reactions is comprehensively elucidated in Chapter 1. Nevertheless, as basis for the following discussion of practical requirements, some fundamental features and definitions of electrolysis will be recapitulated here, sometimes knowingly in a simplified form. If electrochemical knowledge, which is included in this book, has to be transferred into application as a preparative electrolysis in laboratory—or even into industrial electrochemistry—numerous practical aspects have to be considered. Selected issues are briefly discussed in this chapter: various alternatives for electrochemical cell operation, including recommendations for procedure and measurements in electrolysis experiments, an overview of relevant properties of electrode materials, electrolyte components and cell separators, as well as examples of cell constructions using different materials. More detailed information is available in the previous (fourth) edition of this book [1]. The best basis for planning of a preparative electroorganic synthesis is to collect previous information about the desired electrochemical reaction as much as possible. It will be advantageous to know the influence of temperature, solvent, pH value, stirring rate, and so on. Additional to such parameters, which are generally important for chemical reactions, typical electrochemical information, especially the dependence of the processes on the electrode potential, is needed. Electroanalytical standard methods to acquire these data are discussed in Chapter 2. In particular, cyclic voltammetry and coulometry—when necessary combined with other methods—are useful tools for this task. If no special information about the intended reaction is available at the beginning of investigations, also literature data about comparable reactions will be helpful. A comprehensive overview of published electroorganic reactions is given in this book. The desired amount of products determines the scale of the electrolysis cell and the demand of chemicals. Limiting cases may be on the one hand only to prove analytically the formation of compounds or on the other hand to generate a product quantity that is sufficient for application tests. To achieve

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reliable information about an electrolysis process, experiments are suitable on a scale large enough to enable a balancing of reactants, products, and consumed electrical charge (see Section II.F.6).

B. BASIC REQUIREMENTS AND THEORETICAL DEFINITIONS The following characteristics of electrochemistry (and reaction engineering) should be considered for the design of a suitable preparative-scale electrolysis (see also Chapter 1 and literature as, e.g., Reference 2–7). 1. Precondition of Any Electrolysis: A Closed Electric Circuit The electrochemical cell is part of a closed electric circuit. As a consequence, always at least two electrodes have to be used. It is the principle of electrochemistry to replace the direct electron transfer between atoms and molecules in conventional redox reactions by the separated electron release (oxidation) at the anode and the electron consumption (reduction) at the cathode (both processes described with respect to the chemical species in the electrolyte). Nevertheless, in most cases, only one of these reactions may be intended (at the working electrode), but inevitably, the other one has to be carried out (at the counter electrode). Only sometimes useful reactions are possible at both electrodes (so-called paired electrolysis). Even though no valuable product may be achievable at the counter electrode, the reaction there has to proceed at a sufficient rate, and at least any detrimental effect on the desired reaction has to be avoided. Therefore, it is essential to select an optimized combination of electrode materials and a suitable cell: if possible, undivided or divided by a separator if otherwise undesired electrochemical and/or chemical reactions cannot be circumvented (see Section II.E). The transfer of ions between the electrodes is indispensable to realize a closed electric circuit if the electrodes are connected to an electrical power supply, which acts as an electron pump. Thus, an electrolyte of sufficient ion conductivity is needed. Usually, this is a solution containing at least a minimal concentration of an acid, a base, or a salt as supporting electrolyte, which normally has to be separated after the reaction. 2. Types of Electrochemical Reactions Different types of electrochemical reactions within electrolysis processes may be available and the most suitable should be chosen. The practical consequences for cell construction and operation conditions have to be considered (benefits and handicaps of these alternatives are discussed for numerous examples in this book, including industrial applications): • Direct electroorganic electrolysis at an inert or at an electrocatalytically active electrode surface. If this operation mode is possible, the essential properties of electrochemistry are completely applied and no additional agent in the electrolyte is needed. • Indirect electrolysis, using a conventional chemical reaction with an oxidizing or reducing agent that is regenerated electrochemically in a separate electrolysis cell. • Utilization of a redox agent as a mediator, that is, a chemical reaction takes place like in indirect electrolysis, but the mediator is immediately regenerated in situ at the electrode so that it is continuously present like a homogeneous catalyst in the electrolysis cell. 3. Cell Current The cell current within the closed electric circuit (called i in Chapter 1 of this book, usually measured in [mA] or [A]) represents the overall reaction rate of an electrolysis. It induces the charge transfer and the product formation simultaneously at both electrodes, theoretically according to Faraday’s law. The following definitions apply: • Required charge transfer: 1 mol of product needs theoretically 96,485 A s ≈ 26.8 A h (this is the Faraday constant F = charge of 1 mol electrons), multiplied by the stoichiometric number of transferred electrons.

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• Current efficiency (current yield) is a characteristic value of an electrolysis process, being the fraction of the electrical cell current—or (integrated over the time) the fraction of the transferred charge—which is consumed to generate the specified product (separately for anode and cathode). It has to be considered which products result additionally if the current efficiency is less than 100%, for instance, a problematic by-product or a harmless gas. • In order to calculate the required charge for the formation of a product in practical electrolysis, the theoretical value from Faraday’s law has to be divided by the current efficiency. • Current density, called j in this chapter, is the cell current divided by the geometrical electrode area and usually given in [mA cm−2], possibly different for anode and cathode. It represents the local reaction rate of an electrochemical reaction—independent of the size of the entire electrolysis cell—and is a decisive factor for electrochemical reactions due to its close correlation to the electrode potential (see Section I.B.4). Therefore, an even current density distribution on the electrode area is an important demand of cell design. Frequently, a low current density favors the intended electroorganic reaction. However, a higher current density is advantageous considering economic aspects of a sufficient production rate in the electrolysis cell (high space time yield). A compromise has to be found. 4. Electrode Potential The electrode potential—defined in Chapter 1 of this book separately for anode, Φan, and cathode, Φcat—is correlated with the energy conversion during electrochemical reactions, and therefore it is the decisive parameter for thermodynamics as well as for kinetics of any electrolysis. It is defined as the potential difference between the electrode (electron conductor) and the electrolyte (ion conductor) in front of the electrode, which are separated by the electrochemical double layer (see Chapter 1, Section III.A.4). It can be measured in [V] using a reference electrode combined with a (Haber) Luggin capillary (see Section II.F.2 and Chapter 1, Sections II.B.2 and III.A.3). The following elements of the electrode potential have to be considered. a. Equilibrium Potential Reversible electrochemical reactions will spontaneously proceed at an electrode in contact with reactants and electrolyte until the dynamic equilibrium is reached: that is, cathodic reduction and anodic oxidation are running at the same rate and no external current is detectable (the internal bidirectional current flow is called exchange current density = j0). Under these conditions, the electrode will be charged to the characteristic equilibrium potential E, resulting from the charge transfer during the electrochemical reactions. Any electrolysis includes the two equilibrium potentials of the cathode Ecat and the anode Ean. The difference Ecat − Ean is called open-circuit voltage (OCV [V]). It is related to the energy that is at least necessary to perform this electrochemical cell reaction. Multiplied with the correlated transferred charge, the free enthalpy ΔG (Gibbs energy [J mol−1]) of the combined overall cell reaction results. It stands for the thermodynamics of the reaction: ∆G = −OCV ⋅ n ⋅ F

(7.1)

where n is the number of transferred electrons F is the Faraday constant; notice the definition of positive and negative signs The equilibrium potential of the reference reaction H 2 ⇌ 2H + + 2e −

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at standard conditions is set, by definition, to E 0 = ±0 V (NHE = normal hydrogen electrode). The equilibrium potentials E0 of all electrode reactions at standard conditions can be arranged in relation to this value in the electrochemical series. The Nernst equation (see Equation 1.23b) E = E0 +

R ⋅T a ⋅ ln ox n⋅F ared

(7.3)

enables to calculate the equilibrium potential E at real conditions (R = gas constant = 8.314 J K−1 mol−1, T = absolute temperature [K] = [°C] + 273.15, aox, ared = activities of the oxidized and reduced species of the reactants; frequently—mainly at low concentrations—in place of activities approximately concentrations can be used). b. Overpotentials At the equilibrium potential E, by definition, no cell current is flowing. In order to cause a cell current and to start product formation in the electrolysis cell, additionally, overpotentials η have to be applied at both electrodes: A positive overpotential at the anode η = Φan − E > 0 enhances the anodic oxidation and retards the cathodic reduction, and a negative overpotential at the cathode η = Φcat − E < 0 has the inverse effect. Therefore, the overpotentials represent the kinetics of an electrochemical reaction and are the deciding parameters to adjust the reaction rate and generally the performance of an electrolysis. c. Charge Transfer Overpotential (Activation Overpotential) This overpotential is required to overcome the kinetic hindrance of the charge transfer reaction. Its effect has much similarity to the temperature dependency of the reaction rate constant k for chemical reactions, which can be expressed using the empirical Arrhenius equation:  E  k = A ⋅ exp  − A   R ⋅T 

(7.4)

Two fitting parameters are included: A = preexponential factor, theoretically the maximum reaction rate constant at indefinitely high temperature, and EA = activation energy. Only colliding reactants with sufficient thermal energy—according to the Maxwell–Boltzmann distribution—can overcome the activation energy and will react. In the same manner, this model is valid for the charge transfer of electrochemical reactions. However, in electrochemistry—additionally to the thermal energy— electrical energy from the overpotential can be applied for overcoming the activation energy. Thus, in principle, with a sufficiently high electrode potential, each electrochemical reaction could be enforced, even at low temperature (provided that reactions, which can proceed at a lower energy level, are excluded). The empirical Butler–Volmer equation describes—in analogy to the Arrhenius equation—the dependency of the current density j, which represents the electrochemical reaction rate, on the charge transfer overpotential ηCT:  η ⋅α ⋅ n⋅ F   −ηCT ⋅ α cat ⋅ n ⋅ F    j = j0 ⋅ exp  CT an  − exp   R ⋅ T R⋅T      

(7.5)

Two exponential terms with opposite signs consider the anodic and cathodic reactions. The electrical energy, given by ηCT · n · F, is used—dependent on the sign—to decrease or increase the activation energy and in consequence to enhance or retard the anodic or cathodic reactions, respectively. The charge transfer coefficient (symmetry factor) α is the fraction of this electrical energy that influences the respective reaction (αan for the anodic, αcat for the cathodic reaction; for simple reversible

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reactions, α is in the range of 0.5 and αan + αcat = 1). The exchange current density j0 is dependent on the reaction conditions, for example, temperature and concentrations. However, first of all, it is related to the activation energy and hence with the electrocatalytic activity of the electrode for the specified reaction: a low activation energy results in a high exchange current density j0, and in consequence, already a small charge transfer overpotential ηCT is sufficient to generate a major cell current density j (at a high activation energy, the inverse effect occurs). The parameters j0 and α are used for fitting the Butler–Volmer model equation to measured potential current density curves. At elevated positive or negative charge transfer overpotentials ηCT, the term of the respective reverse reaction in the Butler–Volmer equation can be neglected. Then it is possible to take the logarithm (neglecting the identical measurement units of j and j0) and the Tafel equation results (here for a positive overpotential ηCT): ln ( j ) = ln ( j0 ) +

log ( j ) = log ( j0 ) +

ηCT ⋅ α ⋅ n ⋅ F R ⋅T

ηCT ⋅ α ⋅ n ⋅ F ln (10) ⋅ R ⋅ T

[ln (10) ≈ 2.3]

(7.6)

(7.6a)

This is a linear equation of log (j) as a function of ηCT; thus, the current density j increases exponentially with increasing charge transfer overpotential ηCT, for anodic as well as for cathodic reactions (or vice versa, the charge transfer overpotential ηCT increases logarithmically with increasing current density j). The intercept at ηCT = 0, that is, at the equilibrium potential, is log (j0), corresponding to the exchange current density j0. The slope is dependent on the number of transferred electrons n and the charge transfer coefficient α, separately for anode and cathode. A typical value is ≈120 mV additional overpotential per decade of current density (for one electron in the charge transfer step, α = 0.5 and room temperature). d. Concentration Overpotential (Concentration Polarization) In the state of equilibrium, all concentrations at the electrode surface are equal to the concentrations in the electrolyte bulk phase. However, as soon as a cell current is flowing, reactants are consumed at the electrode surface and their concentrations will decrease if their supply is retarded. Analogously, the concentrations of products will increase. The correlating overpotentials result from the Nernst equation as the difference between the equilibrium potential and the potential at the changed concentrations. Typically, two types of concentration overpotential—based on different effects—will occur in electroorganic chemistry: i. Diffusion Overpotential (Diffusion Polarization): Generally, electrochemical reactions take place heterogeneously at the electrode surfaces and mass transfer has a major influence on electrode reactions. Even in a well-mixed electrolyte, there is a stagnant diffusion layer without convection directly adjacent to the electrode, where mass transport of uncharged species is exclusively possible by diffusion. According to Fick’s first law, a concentration difference results, which is proportional to the mass transport and hence to the current density. If this difference is small compared with the entire concentration, the resulting diffusion overpotential, given by the Nernst equation, has only a marginal effect. But with increasing current density, the remaining reactant concentration at the electrode becomes smaller and smaller, resulting in a significant diffusion overpotential. The mass transport of ions (migration) is—in addition to diffusion—enhanced or hindered by the electrical field, dependent on the charge of the ions. The resulting correlation between concentrations and current density can be described in the same way analogously to Fick’s first law. In consequence, all effects of mass transport hindrance result in the diffusion overpotential.

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Finally, if the current density is increased up to the limiting current density, the reactant concentration at the electrode tends to zero. Thus, the maximum possible driving concentration difference is reached, and diffusion cannot be further enhanced. Then, the diffusion overpotential would increase to indefinitely high values. However, in practice, usually a concurrence reaction, like decomposition of the solvent, will start at a high electrode potential (see Section I.B.8). ii. Reaction Overpotential (Reaction Polarization): Decreased reactant concentrations and/or increased product concentrations at the electrode could also be caused by slow chemical reaction steps that are required before or after the electrochemical reaction. In this case— again given by the Nernst equation—a reaction overpotential occurs. 5. Cell Voltage Figure 7.1 elucidates the composition of the cell voltage in an electrolysis cell. Elements are the equilibrium potentials and the overpotentials at the electrodes as described before. Additionally, the ohmic voltage drop (resistance polarization) has to be considered. It is caused by the resistance in the electrodes (electron conductors) and in the electrolytes (ion conductors), including the electrolyte in a cell separator (if applied). The ohmic voltage drop does not directly influence the electrochemical reactions like the electrode potentials. However, it can be a significant part of the cell voltage and influences the following cell characteristics: • Energy demand: Given as the product of cell voltage and cell current. • Heat dissipation: All introduced electrical energy, which exceeds the heat of reaction. ΔH = ΔG + T · ΔS (ΔS = reaction entropy) is lost as heat. Mainly, this is caused by the overpotentials and the ohmic voltage drop. It has to be considered in the context of the temperature control of the electrolysis cell. Local overheating may be possible, for example, in the cell separator.

anode current feeder anode

ohmic voltage drop (electron conductors)

anodic equilibrium potential charge transfer reaction diffusion

anode potential

+

overvoltage anode cell voltage

anolyte (possibly increased by gas bubbles) ohmic voltage drop (ion conductors)

cell separator catholyte (possibly increased by gas bubbles)

overvoltage cathode cathodic equilibrium potential

cathode cathode current feeder

ohmic voltage drop (electron conductors)

cathode potential

charge transfer reaction diffusion

– FIgURE 7.1 Scheme of the composition of the voltage in an electrolysis cell. (From Jörissen, J., in Bard, A.J., Stratmann, M., Eds., Encyclopedia of Electrochemistry, Vol. 8, Ch. 2, p. 35, 2004. Copyright Wiley-VCH Verlag GmbH & Co. KGaA. Reproduced with permission.)

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6. basic Definitions for Chemical Reactions, Also Valid for Electroorganic Electrolysis • Yield, being the fraction of the entire supplied reactant, which has formed the product. • Selectivity, being the fraction only of the converted reactant, which has been used to generate the product. It is a typical criterion to evaluate an (electro-) chemical reaction, especially considering the quality of an (electro-) catalyst. Insufficient selectivity causes losses of reactants and presumably problems in separation of by-products. • Conversion (degree of conversion), being the fraction of a reactant that has been consumed by the reaction. Since the concentrations of reactants are decreased and those of products increased with rising conversion, the selectivity of the desired reaction mostly declines at higher conversion due to diminution of the intended reaction of the reactants and enhancement of consecutive reactions of the products. Typically, a compromise has to be found between increased effort for separation of unutilized reactants at low conversion and intensified by-product formation at high conversion. 7. Operation Modes of (Electrochemical) Reactors Any electrolysis cell—even in a small laboratory scale—is a chemical reactor, and the principles of reaction engineering have to be observed because their influence on the results can be significant. Fundamental differences characterize the two reactor operation modes: • Batch operation, where the electrolysis cell (if necessary anode and cathode compartment separated) is filled with the reactants and operated during a certain time. Subsequently, the products are separated from the solutions. This is the well-known operation mode in chemistry and electrochemistry. Ideally, the entire volume is mixed without any concentration difference (backmix reactor, stirred tank reactor). The conversion and consequently most other conditions in the cell are changing with time. Much information about the electrolysis process is available by analytical monitoring, for example, of concentrations and electrode potentials, during the reaction time (see Section II.F.6). If the reaction conditions change during time—typical for batch operation—it is essential to differentiate between the following: • The actual values of, for example, yield, selectivity, conversion, cell current, current density, charge transfer, and current efficiency. • The summarized (integrated) values from the start to the end of the reaction. • Continuous operation (flow-through cell), where a continuous stream is fed into the cell (if necessary, separated into anode and cathode compartments) and, after partial conversion, exits the cell. Constant operation conditions (stationary state) are reached after a run-in period. Continuous operation is especially interesting for long-time stability tests of cell compounds like electrodes and membranes. For continuous operation, the mixing behavior within the flow-through electrolysis cell is of decisive importance. Two idealized limiting cases can be discriminated: • Backmix reactor with ideal mixing (see item “batch operation”). It is important to consider that the solution in the entire volume is equal to the cell outlet with consequences especially at high conversion (see item “conversion”). • Plug flow reactor, where the flow moves through the cell along the electrode surfaces (theoretically ideal tubular reactor, which means hypothetically no backmixing within the reactor volume but perfect local mixing). Conversion and other operation conditions do not change with time (stationary state) but with location between input and output of the cell.

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• A real flow-through electrolysis cell will be more or less similar to one of these alternatives, dependent on its construction. However, usually the mixing behavior may not be clearly defined, and theoretically calculated results of electrolysis possibly cannot be practically achieved (see cell constructions in Section II.G.3). 8. Control of Electrochemical Reactions The possibility to control the reaction by electrical parameters, which can easily be adjusted— especially using the direct equivalence of cell current and reaction rate—is a typical advantage of electrochemistry in comparison to conventional chemical reactions. Knowledge about the correlation between the potential at the working electrode and the current density is necessary to find a suitable operation mode. Figure 7.2 shows a simplified example for an anodic oxidation reaction, simulated using the elements of the electrode potential that are discussed before. Practically, such a diagram is available, for instance, by cyclic voltammetry or rotating disk experiments with a very low scan rate. The thin line demonstrates the anodic oxidation of the electrolyte (solvent and/or supporting electrolyte) without reactants at an elevated potential, here if the equilibrium potential of this “background” reaction at 0.8 V is exceeded. The current density rises exponentially with increasing charge transfer overpotential. This reaction is undesired at the working electrode and can be neglected here below 0.8 V (for selection of the electrolyte, see Section II.D). This represents the upper limit of the potential window where organic oxidation reactions are possible without decomposition of the electrolyte. If a reactant 1 is added, it can be oxidized according to the thick compact lines. The equilibrium potential is about 0 V, and a significant current density for this desired reaction starts above 0.2 V. However, at increasing potential, soon a limiting current density is attained, which is proportional to the concentration of reactant 1 (indicated for C1 and two-thirds and one-third of it). Even applying a high potential (that means at a high diffusion overpotential), the current density cannot be further 100 Reactant 2 Constant Concentration C2

90

C2 70 C2 Reactant 1 Changing Concentration C1

50 40

decomposi tion

60 C2

C1 30

Electrolyte

Current density/mA cm–2

80

2/3 C1

20

1/3 C1

10 0 0

0.2

0.4 0.6 0.8 Electrode potential/V versus NHE

1

1.2

FIgURE 7.2 Current density–potential curves for the anodic oxidation of two different reactants and finally of the electrolyte. The electrode potentials are related here to the NHE as reference electrode. (From Jörissen, J., in Bard, A.J., Stratmann, M., Eds., Encyclopedia of Electrochemistry, Vol. 8, Ch. 2, p. 35, 2004. Copyright Wiley-VCH Verlag GmbH & Co. KGaA. Reproduced with permission.)

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enhanced and the oxidation runs selectively at the shown constant rate (constant current density) until the potential of electrolyte decomposition is reached. The situation changes if an additional reactant 2 is present that can be oxidized above 0.4 V (thick dotted lines). In this example, a constant concentration C2 is assumed. The limiting current density for the oxidation of reactant 2 is achieved at about 0.8 V and is significantly higher than for the oxidation of reactant 1. Under the conditions of Figure 7.2, a selective oxidation of reactant 1 will be possible at 0.3 V. Reactant 2 can be oxidized only simultaneously with reactant 1 and the respective selectivities are dependent on the concentrations. Analogous correlations have to be considered for the counter electrode (here the cathode). A  suitable reaction has to be performed at the counter electrode, but undesired reactions should be avoided, if necessary using a divided cell with a separator (see Section II.E). Different control modes of electrolysis cells are possible as follows. a. Potentiostatic Operation (Operation at Constant Electrode Potential) As discussed before, the selective oxidation of reactant 1 in the example of Figure 7.2 will be optimally realized at a constant potential of 0.3 V of the working electrode (here the anode). Then, the maximum possible current density (reaction rate) is automatically adjusted, independent of other parameters. This is especially important for batch operation with changing reactant concentrations. Thus, potentiostatic operation in principle is the optimal control mode from the electrochemical point of view, even though problems are possible (see Section I.B.8.d). A suitable electrode potential–measuring equipment in the cell and a potentiostat as shown in Figure 7.3 are required. Special practical aspects of potential measurement are discussed in Section II.F.2. Potentiostatic operation is relatively expensive, especially in larger-scale cells, which need high power (the required quality and in consequence the costs of the potentiostat should be carefully reflected). For charge balancing, an integrator of the changing cell current is necessary (in case of data recording, this is possible using software). b. Galvanostatic Operation (Operation at Constant Cell Current) Operation of an electrochemical cell at constant cell current is uncomplicated, using an inexpensive power supply (also a potentiostat normally can work in galvanostatic operation), and charge balancing is easy. However, suitable results require selection of a current density where a clear potential difference between desired and undesired reactions exists, for example, less than 30 mA cm−2 for the oxidation of reactant 1 in Figure 7.2. This is realizable especially for continuous operation of flow-through cells (steady state) and is typically applied in industrial electrolysis. For batch operation with changing concentrations, a constant cell current is appropriate only if exclusively harmless side reactions can occur. Otherwise, potentiostatic operation will be preferred. c. Operation at Constant Cell Voltage This may be the simplest electrolysis operation, using a cheap power supply. Nevertheless, the target remains generally to apply the optimal potential at the working electrode. This is difficult to realize at constant cell voltage due to the large number of different parts of the cell voltage that are dependent on various parameters (see Figure 7.1). Therefore, operation at constant cell voltage is unusual. d. Challenges in Controlling of Electroorganic Reactions The discussed correlations between cell current and electrode potentials (overpotentials) are applicable without problems in case of reversible electrochemical reactions (see also Chapter 1). However, particularly in case of organic electrochemistry, the reaction mechanism may be a complex chain of electrochemical and chemical reactions, including energy-rich intermediates and irreversible reaction steps. Then it will be impossible to measure an equilibrium potential.

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Potentiostat

mA

Control input

Direct current source

Luggin capillary

Counter electrode Diaphragm

RE

Reference electrode

Working electrode

mV

FIgURE 7.3 Scheme of an electrolysis in potentiostatic operation. The potentiostat ensures a constant potential at the working electrode. The control circuit of the potentiostat determines this potential using a Luggin capillary, which is connected to a reference electrode (RE, see Section II.F.2.a). It adjusts this potential to the voltage at the control input by regulating the cell current between working and counter electrodes. (From Jörissen, J., in Bard, A.J., Stratmann, M., Eds., Encyclopedia of Electrochemistry, Vol. 8, Ch. 2, p. 35, 2004. Copyright Wiley-VCH Verlag GmbH & Co. KGaA. Reproduced with permission.)

Also a theoretical calculation using thermodynamic data in Equation 7.1 may be scarcely useful for application, because the Gibbs energies ΔG are similar for many organic reactions. Thus, there are only little differences in the theoretical equilibrium potentials, different from Figure 7.2, where significant potential differences are available. In consequence, concurrent reactions and only a poor selectivity of the electrode reactions may be possible. Additionally, an effect similar to a large activation energy is caused if the reaction mechanism includes species that need a high energy, for example, energy-rich radical ions. Then, a major charge transfer overpotential—up to 1 V or more—can be observed, and the small differences in the equilibrium electrode potential become irrelevant. Furthermore, complicated and only partially understood influences may be found, for instance, of the electrode material (possibly including its history) or of the electrolyte composition, including unknown impurities. In consequence, sometimes it can be difficult to reproduce electrolysis results. Nevertheless, usually an electroorganic electrolysis can be successfully operated on the base of empirical and experimental knowledge. However, particularly in this case, the discussed challenges necessitate the careful selection of optimal operation conditions as well as of an appropriate electrochemical cell.

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II. COMPONENTS OF ELECTROORgANIC REACTION SySTEMS An impression of some usual or innovative cell components and materials is given in the following; a more detailed overview is available, for example, in References 1,8,9.

A.

ELECTRODES

The characteristic and most important components of any electrochemical cell are the electrodes— particularly the working electrode. In most cases, they are decisive for the success of an electroorganic synthesis. Electrode materials require a sufficient electronic conductivity and corrosion stability as well as, ideally, a selective electrocatalytic activity that favors the desired reaction due to a low overpotential. Simultaneously, the overpotentials for undesired reactions should be high. Different requirements have to be considered for the working and counter electrodes, respectively. For example, anodic oxygen or cathodic hydrogen evolution, respectively, by decomposition of water as solvent is undesired at the working electrode, and high overpotentials for these reactions are favorable. At the counter electrode, such reactions may be intended, and a low overpotential is beneficial. Typically, the electrode reaction includes several steps, such as adsorption and desorption, a single or multiple electron transfer(s), and preceding and/or subsequent chemical reactions. All these steps, and as a consequence the selectivity of the reactions, will be dependent on the properties of the electrode surface. Typical examples are chemical composition, morphology, and porosity. These all may be additionally influenced by the electrode history, for example, removing of a surface contamination or roughening due to corrosion. Moreover, there will be considerable interdependencies between the electrode properties and the electrolyte composition of reactants, products, solvents, and supporting electrolytes, possibly including impurities. As a consequence of such influencing factors, a more or less long-lasting run-in period may be necessary until reproducible results are achievable. Special problems can be caused due to passivation of the working or counter electrode surfaces by insulating layers. Examples are tight oxide films on metals, formed at a high anodic potential, or polymer deposits, generated by anodic oxidation of olefinic or aromatic compounds. Then, the activity of the electrode is reduced, because the surface is partially blocked, and a decrease in the cell current may be necessary. Examples of remedies are as follows: • Periodical changing of the polarity of the electrodes (a symmetrical construction of the cell should be provided) • Additives for an increased polymer solubility in the electrolyte Rarely, in disadvantageous cases, it may be difficult to observe reproducible results of an organic electrolysis due to unforeseeable circumstances at the electrodes. 1. general Requirements for Electrode Materials Electrodes can be made of the following: • A homogeneous material that is electrocatalytically active itself • A homogeneous material that forms in situ an active layer on the surface • A carrier material with an active coating Additionally to these decisive electrochemical properties, experimental requirements for successful investigations have to be considered: • Leak-proof installation of electrodes in the cell body • Proper connection with low ohmic resistance to the current feeder • Appropriate assembling of the cell

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Moreover, the selection criteria for electrode materials include mechanical and practical properties as follows: • Stability, rigidity, elasticity, brittleness, etc. • Ability to be converted to sheets, wires, grids, expanded metal sheets, and porous plates, such as sintered or foamed material and felt • Possibilities of machining, cutting, welding, or soldering • Corrosion resistance • A possible contamination of electrolyte and products has to be considered. • Especially important in case of toxic materials. • Corrosion determines the electrode lifetime (possibly important expense factor). • Price, which may be deciding, particularly in case of commercial use In Sections II.B and II.C, typical electrode materials for application as anode and/or as cathode, and then electrode designs of practical interest are discussed. A comprehensive overview about electrodes is given, for example, in References 8 and 9. 2. Special Requirements for Cathode Materials A typical requirement for a cathode material, used as working electrode in aqueous media, is a high hydrogen overpotential ηH as condition for a strong reduction power. Different materials include a large range of overpotentials (here at 1 mA cm−2, 25°C, in acidic solution [3,7,10]): • Very low ηH < 0.1 V Pt, platinum metals (also active hydrogenation catalysts), Au • Low ηH ≈ 0.2 V Ni (also active hydrogenation catalyst) • Medium ηH 0.2−0.6 V Fe, Cr, Ag, Al, Ti, Mo, W, Bi, graphite stainless steels (Cr–Ni–Fe), brass (Cu–Zn), Monel® (Cu–Ni) • High ηH ≈ 1.0 V Sn, Zn, Cd, Pb, Hg • Very high glassy carbon, Ebonex® (TinO2n−1, n ≈ 4), Boron-doped diamond (BDD) (very low electrocatalytic activity) In nonaqueous media, the hydrogen overpotential is less important. A specific problem can be that a pure cathode material has a sufficiently high hydrogen overpotential, which, however, may be decreased even by very small amounts of other metals (poisoning), and hydrogen evolution occurs. Possible sources of such metals can be impurities in the electrode metal itself and in all applied chemicals, but frequently also corrosion of an (noble metal) anode. For the selection of the anode material, this effect should be considered. Hydrogenation is another typical electroorganic cathode reaction that requires the choice of a cathode material with specific catalytic activity (see Chapter 44). At the cathode as counter electrode, usually hydrogen is evolved. An undivided cell is applicable, if the hydrogen overpotential and the hydrogenation activity of the cathode material are adequately low in order to avoid undesired reduction and/or hydrogenation reactions. 3. Special Requirements for Anode Materials A high positive potential at the anode favors corrosion so that corrosion resistance is an important property of anode materials (unless metal dissolution is intended using a sacrificial anode, see Section II.C.3.). The stability of an anode material is strongly influenced by the operation conditions and by the anolyte composition, for example, temperature, aqueous or nonaqueous medium, pH value, and presence of halides. In most cases, the electrochemical behavior of anode materials in aqueous electrolytes is dependent on in situ formed surface oxides. Also adsorbed compounds, which are formed by oxidation from reactants—a well-known example is carbon monoxide—may be of decisive influence.

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The oxygen overpotential is a significant attribute of an anode material for use in aqueous electrolytes. It should be high for the anode as working electrode, especially if a strong oxidation power is required, in order to minimize concurrent oxygen generation. During oxygen evolution at the anode as counter electrode, undesired oxidation reactions are hardly avoidable at the high oxygen potential, including the oxygen overpotential. A low-oxygen overpotential would be helpful; however, this property is not available using known materials. Therefore, frequently, a divided cell is required where the anode as counter electrode operates in a separate compartment (a sacrificial or gas depolarized anode could be an alternative, see Sections II.C.3 and II.C.4).

B.

EXAMPLES OF ELECTRODE MATERIALS

1. Platinum, Platinum Metals, Other Noble Metals, and Their Alloys These metals are classical electrode materials that are very often applied in the literature. The main reasons are the high electrochemically as well as chemically catalytic activity and the resistance to corrosion in nearly all solutions. These properties show significant differences between the various metals and alloys of this group so that a careful selection is necessary. A considerable influence of the orientation of single crystals is observable; nevertheless, only polycrystalline materials will be appropriate for synthesis applications. It can be economically interesting to use these very expensive metals as a thin coating on a carrier (see Section II.B.8). a. Cathode Platinum and palladium—and comparably the other platinum metals—operate with the lowest known overpotentials for hydrogen. Simultaneously, they are the most effective catalysts for hydrogenation reactions [11]. b. Anode The anodic behavior of these metals is strongly influenced by surface oxides and/or adsorbed compounds (this can be demonstrated by cyclic voltammetry, see Chapter 2). For example, the anodic activity of platinum is intensely declined by adsorbed carbon monoxide, while the oxidation of carbon monoxide to carbon dioxide is enhanced using a mixed ruthenium platinum catalyst (this is important, e.g., in the direct methanol fuel cell [DMFC] [12]). Usually, corrosion is no significant problem for platinum metals under most conditions and anolyte compositions (probably, it has to be considered economically for commercial applications). Nevertheless, as previously mentioned, even traces of anode corrosion can cause a decrease in the necessary high hydrogen overpotential at the cathode. Anodic corrosion of this group of metals is very dependent on the anolyte pH value, particularly in case of the less noble metals. 2. Nickel a. Cathode Nickel has a relatively low hydrogen overpotential as well as a high hydrogenation activity and may be an inexpensive alternative for platinum metals, as working and also as counter electrode. Its corrosion resistance as a nonnoble metal has to be checked for the actual conditions. Typically, nickel is suitable for aqueous alkaline solutions. Very fine dispersed Raney nickel offers an increased activity. It is prepared from a layer of a nickel alloy with aluminum or zinc on the cathode surface. Prior to use, the alloy metal is dissolved in alkaline solution (see, e.g., Reference 13. Caution: Raney nickel is attacked by oxygen and self-ignition in air is possible). b. Anode (Active Surface: NiOOH) A layer of NiOOH is formed on the surface of a nickel anode in aqueous alkaline solution. This is intensively studied [14] because it is used in nickel cadmium and nickel metal hydride

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accumulators. In electroorganic chemistry, selective oxidation reactions are possible [15]. Additionally, it can be used as oxygen-evolving counter electrode. 3. Iron (Mild Steel), Stainless Steel a. Cathode Iron has a relatively low hydrogen overpotential. Nevertheless, selective cathodic reactions are published, for example, Reference 16. Due to its very low price, it is favored for commercial applications, mainly as counter electrode. Corrosion can occur when the cell current is switched off. An alternative, particularly for laboratory cells, may be stainless steel with much better corrosion resistance and only marginally increased hydrogen overpotential. 4. Lead The classical electrode material lead is largely used in lead–acid batteries. Therefore, it has been comprehensively investigated in the literature, for example, Reference 17. The mechanical stability of the very soft pure lead is insufficient for application. Alternatives are lead-coated carrier electrodes or tougher alloys (typically with antimony: type metal). a. Cathode Lead cathodes achieve a strong reduction activity in aqueous solutions due to the very high hydrogen overpotential. However, this requires very pure lead and a clean procedure. If the cathode reaction has to operate at the limit of hydrogen evolution, the presence of other metals with a lower hydrogen overpotential is detrimental (poisoning). Sources, which have to be carefully excluded, may be impurities or alloying metals in the lead, contaminations of the electrolyte and the reactants, and possibly corrosion of the anode [3]. b. Anode (Active Surface: PbO2) The surface of lead as anode material is oxidized at high anodic potentials in acidic aqueous solutions to lead dioxide (e.g., Reference 18). The continuous growing of this layer destroys the lead base metal during longer operation. The overpotential for oxygen evolution on PbO2 in acidic aqueous solutions (frequently sulfuric acid) is very high so that a high anodic potential and a strong oxidation activity are enabled. However, PbO2 simultaneously is a powerful chemical oxidation agent, and the electrode may be damaged by a spontaneous reaction with a reactant. Basically, the stability of pure PbO2 anodes is relatively poor. Reasons can be dissolution in the electrolyte and frequent mechanical erosion. Considering the toxicity of lead, this can be a problem. Meanwhile, high stability has been achieved by the development of PbO2 coatings on suitable carrier materials (see Section II.B.8) with sophisticated additives (comprehensive review [19]). PbO2-coated titanium anodes are commercially available (e.g., from References 20 and 21). 5. Mercury a. Cathode Mercury is traditionally the most used cathode metal for electroorganic reduction reactions. In aqueous solutions, it has the highest hydrogen overpotential of all metals. Another property is its liquid state, which has to be considered for cell constructions (see Section II.G.2). This enables— simply using a (magnetic) stirrer—a continuous renewing and cleaning of the cathode surface. At very negative potentials, all present metal cations can be reduced, possibly even alkali metals of a supporting electrolyte. However, these metals are dissolved as amalgam and the risk of decreasing the hydrogen overpotential (poisoning) is less than at a solid metal surface. Dependent on the application, an amalgamated electrode of copper or another metal may be appropriate for easier handling.

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The application of mercury is limited due to its toxicity, and careful safety precautions are required. The hazard of mercury contamination restricts its use for commercial syntheses. 6. Carbon Carbon has several advantages as less expensive electrode material that is available in very different variations (e.g., References 22 and 23). Usually, its electrocatalytic activity is relatively low. It is used in large amounts for inorganic and organic industrial electrolysis as well as for batteries. Appropriate constructions of electrodes and current feeders enable the application of carbon despite its low conductivity (about one-hundredth of most metals). Traditional carbon materials for electrode applications differ in the content of crystalline graphite, conductivity, and mechanical and chemical stability. Usually, they are impregnated with chemically resistant resins in order to reduce the originally high porosity and become leak-proof against gases and fluids. Such materials are brittle, their solidity is limited, and their machinability is good. Flexible sheets of graphite are produced as sealing material (e.g., SIGRAFLEX®, SGL Group). It is easy to handle as a corrosion-resistant electrode or current feeder; however, during gas evolution, swelling and erosion can occur. Fiber materials of carbon or graphite are offered, for example, as tissue, felt, or paper. Typically, they are used as conductive diffusion zones in fuel cells. Additionally, they are interesting electrode materials of very high porosity (free space volume up to 80%, 3D electrodes) and with remarkably good elasticity. Glassy carbon (vitreous carbon) is another carbon modification (ceramic-like, based on microstructures of fullerene instead of graphite [24]). It is an interesting electrode material, if very high corrosion resistance is required. Additionally, its electrocatalytically activity is low; thus, a large potential window in aqueous solution is accessible. Methods of surface modification and doping are reviewed in Reference 25. Glassy carbon is rigid, smooth, and tight against gases and liquids. The price is relatively high but lower than for noble metals. It is available in simple shapes and as a foamed material (reticulated vitreous carbon for 3D electrodes [26,27]). Due to its hardness, machining is restricted (only with diamond tools). Further carbon-based electrode materials are carbon-filled polymers or a carbon paste with paraffin, which is interesting for laboratory investigations, because a surface layer can easily be removed for regeneration. a. Cathode Carbon has a relatively high hydrogen overpotential so that manifold reduction reactions are possible, if the cathode is the working electrode. Additionally, the catalytic activity for hydrogenation reactions is low. This is beneficial, if a carbon cathode is the counter electrode: besides the intended hydrogen evolution, undesired cathodic hydrogenation reactions are prevented (such reactions, in the worst case, could attack the products of the anode). b. Anode Carbon anodes are destroyed by oxidation to carbon dioxide, if oxygen is evolved in aqueous media. This is enhanced with increasing porosity, while glassy carbon is relatively stable. The oxygen evolution can be hindered at a lower pH value; however, generally, carbon anodes may not be optimal in aqueous solutions. Moreover, intercalation of anions can decrease the stability of carbon anodes. Carbon (graphite) is a typical anode material in nonaqueous media. For example, it is suitable for methoxylation reactions, even in an industrial scale, especially due to its increased overpotential for the oxidation of methanol. 7. Conductive Ceramic, for Example, Ebonex As an alternative for glassy carbon (see Section II.B.6), Ebonex can be an interesting electrode material, if high corrosion resistance, coupled with low electrocatalytic activity and high overpotentials for oxygen and hydrogen, is required [28]. Furthermore, it is a suitable carrier material for electrode

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coatings (see Section II.B.8). Ebonex is made of Magneli phases, substoichiometric titanium oxides with reduced oxygen content TinO2n–1 (3 < n < 10, mainly n ≈ 4), which are sufficiently electrically conductive. As a ceramic material, it is hard and brittle with limited possibilities of machining (diamond tools). 8. Electrocatalytic Coatings on Carrier Materials Thin layers of noble metals on a less expensive carrier material are frequently used as cost-efficient alternative for compact noble metals. Additionally, there are several interesting electrode materials published in the literature that can be applied only as a coating. For example, some transition metal oxides enable a considerably high reaction selectivity. However, their relevance in electroorganic synthesis until now is low due to their insufficient stability (no detailed discussion is provided here). An up-to-date exception is the BDD electrode, whose development meanwhile achieved excellent durability (see Section II.B.9). If a corrosion-resistant and electrochemically less active carrier material is used—for example, glassy carbon or Ebonex—even very thin layers of a coating will be sufficient for laboratory smallscale experiments with low current densities. Then, the electrochemical properties of the electrode are correlated only with the coating but not with the carrier (e.g., Reference 29). Titanium metal is a typical carrier metal for larger electrodes up to an industrial scale, especially for inorganic electrolysis. In aqueous electrolytes under oxidizing conditions (even in the presence of chlorine), it is excellently corrosion resistant due to a stable, self-healing TiO2 passivation layer. However, under reducing—and especially acidic—conditions, titanium is unstable, intensified in the presence of organic acids or fluorides. Usually, for nonaqueous organic media, it is inapplicable. The oxide layer on titanium is semiconducting and inhibits an electron transfer in anodic direction (valve metal). In principle, an application as cathode may be possible, but if hydrogen is evolved, titanium is destroyed by the formation of titanium hydride. Platinum-coated titanium electrodes as replacement for pure platinum are commercially available (e.g., from References 20,21 and 30). For sufficient stability, the platinum layer has to be porefree tight. Lead dioxide–coated titanium electrodes are mentioned in Section II.B.4. Dimensionally stable anodes (DSA®) are the most used anodes in industry [31]. Here, the TiO2 passivation layer on the titanium is replaced by a conductive coating, based on titanium and ruthenium oxides, with optimal electrocatalytic activity for chlorine evolution from aqueous chloride solutions. Comparable coatings, mainly based on iridium and tantalum oxides, are available for oxygen and for simultaneous chlorine and oxygen evolution. Such titanium anodes also may be interesting for electroorganic syntheses, for example, as oxygen-evolving counter electrode and probably as working electrode (possible suppliers are, e.g., References 20,21 and 30, which is focused on special anodes). However, previously, the stability of such electrodes in the presence of the applied organic compounds has to be checked. 9. boron-Doped Diamond Coating A relative new electrode material with outstanding properties and excellent chemical durability is diamond, which obtains electrical conductivity by doping with boron (e.g., Reference 32). It is coated onto various high melting carriers—for example, silicon, carbon, tantalum, niobium, titanium, conductive ceramics—by chemical vapor deposition. Its catalytic activity is low, and in aqueous solution, it achieves the highest known overpotential for oxygen and also a high one for hydrogen evolution. Thus, it offers an extraordinarily high potential range and electrochemical power for oxidation as well as for reduction. These properties of BDD electrodes appeared useful for mechanistic studies [33–35], spectroelectrochemistry [36], as well as preparative applications up to industrial scale [37], which have recently been reviewed [38]. The high accessible positive potentials allow “incineration” [38] and destructive detoxification by extremely reactive radical species [38], which may, on the other hand, turn out to be problematic for the selectivity of syntheses.

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281

EXAMPLES OF ELECTRODE TYPES AND THEIR SPECIAL PROPERTIES

1. Porous vs. Smooth Electrodes The active area of a porous electrode can be much larger—up to about three orders of magnitude—compared with the geometrical area that is active at a smooth electrode. Thus, the effective current density is substantially reduced with significant influence on the electrochemical behavior at a decreased sum of overpotentials. However, there is a diffusion overpotential included due to mass-transfer effects, which are dependent on the porosity and also on the shape, dimension, and structure of the pores within the electrode material (a wide range of variations is possible). Classical examples for increased electrode surface are platinized platinum (electrochemically deposited, e.g., Reference 39) and Raney nickel (see Section II.B.2). Here, reactions are enabled that are impossible at smooth electrodes of the same metal. The electrode activity can also be enhanced by a simple roughening of the surface, for example, mechanically or by etching. Three-dimensional electrodes are a direct way to expand the effective electrode area. Typical examples are carbon fiber felt, sintered, foamed, or reticulated materials, or a packed bed of particles. Special attention is required to include a significant depth into the working area, as part of the porous electrode thickness, more than a small layer at the surface (the conductivity ratio of the electrode material and the electrolyte is important; see, e.g., References 40 and 41). 2. gas-Evolving Electrodes Frequently, gases are evolved—at the working as well as at the counter electrode—either during the main reaction or as an undesired by-product (e.g., carbon dioxide, hydrogen, oxygen, chlorine). Caution: safety precautions (at least a splinter shield) are indispensable, if explosive gas mixtures and an ignition by a shortcut of the electrodes are possible. Usually, the movement of gas bubbles is advantageous because it enhances the desired mass transport at the electrode surface. However, gas bubbles increase the voltage drop between the electrodes, and a suitable electrode design is necessary to remove gases quickly from the front to the backside of the electrode, for example, using a mesh or an expanded metal sheet. An additional problem of (big) gas bubbles is a deformation of the even current density distribution at the electrodes and, in consequence, a possible degradation of the reaction selectivity (see Section II.F.1). In the worst case, the electrode surface could be partially disabled. Thus, if there is an intense gas evolution, the removal of gas from the gap between the electrodes and out of the electrolysis cell is a significant cell design attribute (see Section II.G, theoretical discussion in Reference 42). 3. Sacrificial Anodes Usually, a sufficient corrosion resistance is required for electrodes, especially for anodes. However, in case of a sacrificial anode, its disintegration is the desired reaction. It is a special type of depolarized electrode where a reactant is delivered—here the anode material itself—for an additional electrode reaction [43]. The sacrificial anode will be the working electrode, if the synthesis of an organometallic compound of the anode material is intended (see Chapter 35). Another example is polysulfides, which are producible from sulfur as anode material, which is mixed with carbon powder for adequate conductivity [44]. Typically, sacrificial anodes are applied as counter electrodes. Using a strongly electronegative metal—for example, zinc, aluminum, magnesium—no anodic oxidation is possible at the very negative potential, and an undivided cell is utilizable. This technique can be operated also in nonaqueous media ([43], industrial example [45]). Anions have to be available as counterions for the produced metal cations (generated either by the cathodic reaction at the working electrode or added as an acid whose H+ ions are consumed there). It will be useful, if the produced salt is soluble enough for sufficient conductivity of the electrolyte, but its excess should precipitate for easy separation.

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4. gas Diffusion Electrodes Depolarized electrodes of another type apply a gas as additional reactant. However, the electrochemical conversion of a gas is much more complicated than gas evolution. Gas diffusion electrodes (GDEs) are necessary: the gas, the electrolyte, and the catalyst—electrically connected with the current feeder—have to be in optimal contact at a maximized electrode surface (three phase zones within a highly porous electrode structure). GDEs are well known for fuel cells and object of intensive research (overview, e.g., References 46 and 47). GDEs may be interesting also for electroorganic synthesis (until now, only for aqueous electrolytes), though actually there are merely a few results published (e.g., References 48 and 49). Nevertheless, a DMFC—in principle—is an electroorganic reactor for the oxidation of methanol in aqueous solution in combination with an oxygen-depolarized cathode (ODC) [12,50]. GDEs could be advantageous especially as counter electrodes. The potential of an oxygenconsuming cathode (= ODC) is significantly more positive than a hydrogen-evolving cathode under the same conditions (theoretically shifted by 1.23 V, practically by up to 1 V). Thus, undesired reduction reactions are nearly impossible. Vice versa, a comparable shift in negative direction is achieved by a hydrogen-consuming anode (possible alternative for a metallic sacrificial anode, see Section II.C.3). Here, unwanted oxidation reactions are excluded (e.g., no chlorine evolution in the presence of chloride ions).

D.

ELECTROLYTES

The typical liquid reaction medium for organic electrolyses consists of a solvent and a supporting electrolyte. This solution is commonly called the electrolyte. It does not only provide the environment for the electrochemical reaction but also ensures the ionic charge transport in the cell between the working and counter electrodes. 1. Solvents An excellent review of solvents used in organic electrochemistry forms part of Lund’s chapter on Practical Problems in Electrolysis in the fourth edition of this monograph [1]. The respective section is reproduced as Appendix A. Izutsu [51] provides another collection. 2. Supporting Electrolytes Similar to Section II.D.1, the commonly used supporting electrolytes for electroorganic applications have been described extensively by Lund [1]. The section dealing with this topic is reproduced as Appendix B. Again, Izutsu’s monograph [51] gives additional information. A noteworthy development, which has not yet been foreseeable in the fourth edition, is the increasingly popular use of ionic liquids in organic electrochemistry. Chapter 8 of the present edition is devoted to this evolving field, as well as to electrochemistry in emulsion-based and supercritical fluid (SCF) electrolytes. Here, we discuss some additional aspects of supporting electrolytes. Electrolytes on the basis of weakly coordinating anions have recently been described as very useful for mechanistic investigations [52–57], in particular with respect to the separation of formal potentials in multielectron transfer reactions (see Chapter 11). Such anions are often highly fluorinated [52,53] or derivatives of closo-borates [57] and are characterized by a large volume over which the negative charge is distributed [53]. As a result, a low tendency toward ion pairing with cations is observed and shielding of the positive charge is decreased [54]. Furthermore, the nucleophilicity of such anions is weak [53], minimizing side reactions of electrochemically generated electrophiles. Another advantage of such electrolytes is a more extended potential window in the positive potential range [52]. The popular quaternary ammonium cations of supporting electrolytes are well known to undergo Hofmann elimination upon reduction (see Appendix B, Section III.B.2.d). Recently, another side reaction during oxidation was identified that produces alcohols, ketones, and amides [58].

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The need to separate the supporting electrolyte after electrolysis causes additional workup efforts. Electrosyntheses in the absence of such salts would simplify the overall processes. Apart from solid–polymer–electrolyte (SPE) technology (see Section II.D.3), electrochemical experiments in organic solvents in the absence of supporting electrolytes have been reported as early as 1984 [59]. Such conditions were used mainly in mechanistic work at ultramicroelectrodes [60], which is then complicated by migration effects [61]. However, recently, electrosynthesis in organic solvents without supporting electrolyte was also described in miniaturized environments (see Chapter 9) [62]. 3. Solid Polymer Electrolyte Technology Generally, an ion exchange membrane is working as an ion conductor if it is used as a cell separator (see Section II.E.2). Thus, it is also able to function directly as an electrolyte even in the absence of an ion-conducting solution. This SPE technology is established in fuel cells (proton exchange membrane fuel cell [PEMFC] and DMFC) [12,46,47,50]. Analogously, electroorganic syntheses without any supporting electrolyte are possible in comparable electrolysis cells if a sufficient conductivity of the membrane can be achieved in the reaction system (see Section II.E.2). In this case, operation is easy, and expenses for separating and recycling a supporting electrolyte are economized [63]. For various reactions applying cation and anion exchange membranes, combined with different electrode materials in aqueous and even in nonaqueous media, possible advantages of the SPE technology have been demonstrated [64–66].

E.

CELL SEPARATORS

The easy construction and operation of undivided electrochemical cells is generally advantageous (see Section I.B.1). However, if the success of the electrolysis is impeded by an unavoidable reaction at the counter electrode, resulting, for example, in significant deficits of yield and/or contamination of the products, a separator between anode and cathode will be necessary. No ideal separator is available; always a compromise of several properties has to be found. On the one hand, the characteristic of a cell separator is to prevent a direct mixing of anolyte and catholyte and to decrease diffusion in both directions; on the other hand, the voltage drop during the obligatory migration of ions has to be minimized. A possible overheating as a consequence of an excessively high voltage drop should be considered. With increasing differences in the properties of anolyte and catholyte, the functioning of the separator becomes more and more difficult. Diffusion is intensified by increasing concentration gradients. A less porous and/or thicker separator will diminish diffusion but increase the voltage drop. Precipitation of a compound at the surface or within the separator in case of different solubilities in both electrolytes can be a problem and must be prevented. A nearly optimal separation of anolyte and catholyte is possible using two separators in series, if in laboratory experiments a high voltage is acceptable. Additionally, the volume between the separators can be rinsed with a suitable solution in order to remove compounds that enter through the separators but are not tolerable in the opposite cell compartment. 1. Porous Materials Characteristics of a porous separator (diaphragm) are pore diameter, porosity, and thickness. Additional aspects are essential for practical application such as mechanical strength (brittle or flexible), chemical stability, and constant dimensions (possible swelling in the solvent). The classical porous separator in laboratory cells is sintered (fritted) glass (frit). A glassblower easily can mount it by melting into the walls of glass cells (leak-proof). Types of different properties are available and also similar sintered ceramic materials are used (e.g., alumina, unglazed porcelain, or pottery). Glass and certain ceramic materials will be attacked by strongly alkaline and acidic fluoride-containing media. In principle, separators of glass or ceramic materials are rigid and

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brittle and need an increased thickness for sufficient mechanical stability in larger dimensions, thus causing a high voltage drop. Glass or ceramic surfaces adsorb OH− ions in aqueous solutions. Therefore, they are negatively polarized and a positive charge appears in a thin layer of the solution close to the surface (ζ-potential). As a consequence, this solution can be moved in the electric field toward the cathode. This so-called electroosmotic transport occurs especially in separators with small pores and therefore large surface. It can produce significant undesired transportation through the separator. Flexible porous diaphragms (down to thin foils) can be manufactured from fibrous materials, for example, as paper, felt, or woven fabric, and are applicable up to large dimensions. Even simple materials such as filter paper, regenerated cellulose film (cellophane), or an agar–agar plug may be useful in case of nonaggressive solutions. Diaphragms made of asbestos exhibit very good technical properties; however, its application as carcinogenic substance today is obsolete. Several polymers are used as base materials for porous diaphragms (e.g., as separators in batteries). Organic solvents can attack polymers, typically by swelling that may clog a diaphragm. The stability of a polymer is strongly influenced by all compounds in the electrolyte mixture. Therefore, a careful selection of the diaphragm in combination with the electrolyte is required, especially in case of long-term applications. Polyethylene and polypropylene are inexpensive and stable at various conditions. Nearly universal stability is assured with porous (expanded) polytetrafluoroethylene (PTFE, e.g., Teflon®, a traditional trade name of expanded PTFE is GORE-TEX®). The hydrophobic properties of such polymers inhibit an immediate application in aqueous media. However, after initial wetting with a completely water-miscible alcohol (e.g., 2-propanol), aqueous electrolytes become applicable. 2. Ion Exchange Membranes While porous separators are permeable for all compounds and ions, there is a transport selectivity for either cations or anions in ion exchange membranes. A large variety is commercially available, especially for utilization in electrodialysis. Thus, their application as cell separator can be interesting, but special properties have to be considered [67]. Ion exchangers and ion exchange membranes consist of a base polymer (typically polystyrene) with covalently bonded fixed ions, typically sulfonate anions or (quaternary) ammonium cations, respectively. For electrical neutrality, counter ions of the opposite charge have to be present. These are cations in a cation exchange membrane (e.g., H+, Na+) or anions in an anion exchange membrane (e.g., Cl−), respectively. They become mobile and exchangeable as charge carrier if the membrane is swollen in a suitably polar solvent, and the fixed and counter ions are sufficiently dissociated and separately solvated. The charge carrier ion in the ion exchange membrane has to be present also in the cell electrolytes. The typical solvent is water that enables a high ion conductivity of ion exchange membranes. Other solvents with adequately high polarity or solvents in mixtures may be possible (experimental checking will be necessary). The material of the ion exchange membrane has to be durable in the cell electrolytes, both chemically and resistant against excessive swelling. A well-known cation exchange membrane of nearly universal chemical stability is Nafion® (Dupont, perfluorinated polymer with sulfonic acid groups). However, it is expensive and its swelling in organic solvents can be prohibitive. A suitable membrane—cation or anion exchange membrane—for a planned application should be selected in contact with the technical support of suppliers. Usually, the transport selectivity of ions will be satisfactory. However, dependent on the conditions—especially due to significant concentration differences—also permeation of ions with the opposite charge and diffusion of other compounds are possible. A characteristic, unavoidable transport mechanism of ion exchange membranes is electroosmosis (enhanced in comparison with porous glass or ceramic; see Section II.E.1). It is generated by the

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solvation shells of the charge carrier ions, because predominantly ions of only one type—cations or anions, respectively—are migrating in a definite direction through the membrane. Unselectively, all solvents, reactants, and products are included in the electroosmotic flow. The magnitude of this effect is influenced by the membrane type and by the composition and concentration of the electrolyte solutions. Up to several molecules per migrating ion can be involved, resulting in a sizable flow. Therefore, this effect has to be considered if an ion exchange membrane will be used as a cell separator.

F. ELECTROCHEMICAL CELLS Optimized electrochemical cells are needed to execute electroorganic electrolysis, considering—as far as possible—all aspects discussed in this chapter. The requirements of the following checklist commonly will be essential; however—dependent on the reaction system and on the goal of the planned research—their significance may be different: • General properties • Suitable materials of electrodes and cell components − Chemically and mechanically stable and not corroding − Optimal electrocatalytic enhancement of the desired reaction and suppression of undesired side reactions (for electrode materials, see Sections II.A and II.B) − Uncomplicated and inexpensive manufacturing (for cell construction materials, see Section II.F.5) • Homogeneous current density on the entire area of the working electrode, that is, equal current distribution at uniform electrode potential • Reliable electrode potential measurement • Constant mixing for steady mass transfer • Well-defined temperature • Effective handling during experiments • Easy and leak-proof mounting • Accurate sampling • Reliable mass and charge balancing • Requirements in special cases may be, for example, • Inert gas atmosphere • Reflux condenser to prevent a loss of solvent (in case of gas evolution, an additional low-temperature cold trap may be necessary) • Balancing of gas evolution 1. Homogeneous Current Density The close interdependence between current density and electrode potential and in consequence the influence on the electrode reactions and their selectivity has been discussed in Section I.B.8. Hence, usually an even current density on the entire electrode area is essential for authentic results. This requires for every point of the electrode area a sufficiently constant overall cell resistance: of the electrodes, of the electrolytes, and—where used—of the cell separator. Satisfactory conductivities of the electrode materials and adequately dimensioned, symmetrical current feeders are required. Only parallel mounted electrodes in a matched cell volume enable a constant resistance of the electrolyte, and in addition, locally invariable concentrations by sufficient mixing are necessary. This is demonstrated in Figure 7.4(a), where the parallel and equidistant lines illustrate the homogeneous current density distribution.

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C C

i

RE

W

C

W

W

L

i

1

i

i

i

i

(a) (b)

(c) RE

RE C

L

RE C

W

i

i

i

L

C

W

i

W

L

i

i

d

∆U = i · R

2d (d) 2/3d

(f)

(e) C RE C

L

W

RE C

W

W

L

i i

i

i

i

i

RE (g)

(h)

(i)

FIgURE 7.4 Scheme of current density distribution and Luggin capillary positions in electrolysis cells. W, working electrode; C, counter electrode; i, cell current; L, (Haber) Luggin capillary; RE, reference electrode; ΔU, resistive voltage drop; R, Ohmic electrolyte resistance.

Inhomogeneity occurs if the cell volume is bigger than necessary for the electrodes, as in the example of Figure 7.4(b). The cell current will use also the electrolyte outside of the electrode shape, where the resistance altogether is smaller and the current density is enhanced at the electrode edges. Inhomogeneity is increased if electrolyte also is present behind the electrodes (Figure 7.4(c)) and some current can flow to the back sides of the electrodes. Additionally, the current distribution will be deformed in case of unsymmetrical cell constructions. The preconditions for homogeneous current density distribution according to Figure 7.4(a) become more and more important in case of increasing current density, decreasing electrolyte conductivity, and reduced electrode distance. Examples of cell constructions, which enable a homogeneous current density, will be discussed in Section II.G. A special case of disturbed current density distribution can be caused by gas evolution, in particular in the upper part of a cell with vertical electrodes. Therefore, a fast displacement of gases from the electrolyte between the electrodes or between the separator and the electrodes has to be provided by electrode and cell construction. If required, the current density has to be limited so that the gases can be released without problems.

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2. Electrode Potential Measurement A precondition of electrolysis at an appropriate potential of the working electrode is a sufficiently reliable potential measurement. It is indispensable for potentiostatic operation (see Section I.B.8). Although it may not be necessary for galvanostatic operation also, in this case it provides important information about the reaction. Generally, potential measurement requires a (Haber) Luggin capillary (see Section II.F.2.c), which connects the electrolyte in front of the working electrode with the electrolyte in the reference electrode. The Luggin capillary functions as a salt bridge that is filled with a fitting electrolyte; a typical example is saturated aqueous potassium chloride solution. The reference electrode has to guarantee a well-defined and well-known potential between its electrolyte and its electrical connector. Exact electrode potential measurement is complicated, especially in case of organic electrochemistry (see also Section I.B.8.d). Possible errors can be minimized only using highly sophisticated methods (see, e.g., Reference 68). However, practical preparative organic electrolysis does not need necessarily the knowledge of exact electrode potentials. Adequately reproducible current density/ potential curves (see Figure 7.2), which are available by electroanalytical methods (see Chapter 2) and should be sufficiently valid equally in electrolysis cells, will be satisfactory. A constant deviation from the exact value in both measurements causes no detrimental effect. Nevertheless, the following practical recommendations should be considered in order to avoid systematic errors, which could have a negative influence on the results of electrolysis experiments. a. Reference Electrodes The reference electrode [69] defines the zero of the potential scale in the electrolysis experiment. For aqueous electrolytes, a variety of (even commercially available) reference electrodes of the second kind (i.e., the potential-determining ion concentration is fixed by the presence of a sparingly soluble salt, e.g., calomel or Ag/AgCl electrodes) is in common use [69a,b] (see also Chapter 1, Section II.B.2). The situation is much more complex in nonaqueous solvents [51,69c]. The common aqueous reference electrodes are often unstable in such environments [70], for example, owing to disproportionation [71]. Separating the aqueous reference electrode from the nonaqueous electrolyte by a diaphragm (e.g., frit, agar–agar stopper) and a salt bridge induces liquid junction potentials [72]. For porous glass plugs, electrostatic ion transfer screening [73] has been described, which is sample dependent and leads to errors of several tens of mV in the reference potential. In addition, contamination in both directions, for example, transport of water into the cell electrolyte, might be a problem. As a consequence, nonaqueous Ag/Ag+ electrodes (silver wire in a solution of a silver salt soluble in the respective solvent, often AgClO4, with a specified concentration and separated from the electrolyte by a frit and a salt bridge) are popular [69c]. Although being reference electrodes of the first kind, for practical purposes, they appear to be stable enough with respect to reproducibility. Their use has been commented critically [74a]. A miniaturized construction for application in ionic liquids has been described [74b]. It is not uncommon, especially in miniaturized cells, to employ chloridized Ag wires [75] or even Pt wires [76] as pseudo-reference electrodes [69c]. The potential of such an electrode, however, may strongly depend on the composition of the electrolyte surrounding the wire and possibly changes in an unpredictable way during experiments. This might be particularly dangerous in solvents containing traces of halide ions (e.g., dichloromethane). Owing to the often high resistance in organic electrolytes, the construction of reference electrodes has received attention with respect to the minimization of ohmic drop and oscillatory artifacts caused, for example, by the reference electrodes’ internal resistance [77a] (see also Section II.F.2.c). Dual reference electrodes have been recommended [77].

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b. Reference Redox Systems in Organic Electrolytes To overcome some of the problems described in Section II.F.2.a, reference redox systems are used. 0 These are redox couples with a more or less reproducible formal potential Eref , which defines the zero of the potential scale. An IUPAC recommendation is the use of the ferrocene/ferricenium (fc/fc+) or bis(biphenyl)chromium (0/I) redox couples [78a]. It has been argued that these redox systems do even present a 0 , as far as possible (extrathermodynamic assumption [78b], see, however, solvent-independent Eref [78c]). Thus, E 0(fc/fc+) can easily be obtained in the respective electrolyte from cyclic voltammetric peak potentials Epox and Epred (see Chapter 2) as the mean (Epox + Epred ) / 2. This value is given vs. a certain reference electrode, for example, a frit-separated Ag/Ag+ electrode with silver ions in a defined concentration, typically 0.01 M, in CH3CN. Note that the resulting value is not necessarily independent of the electrolyte composition owing to, among others, diffusion potentials developing in the frit. It can, however, be used as a reference for any other potential measured in the same electrolyte vs. the same reference electrode:

(

)

(

)

(

E vs. fc/fc + = E vs. Ag/Ag + − E 0 fc/fc + vs. Ag/Ag +

)

The resulting E (vs. fc/fc+) are then comparable. This has been used for the characterization of solvent exchange equilibria coupled to an electron transfer by analysis of E0 shift data in solvent mixtures of a broad composition [79] in order to eliminate artifacts introduced by the solvent variation. When reporting such potential results, it is essential to clearly state the conditions of the reference experiment and how the conversion of the potential values is being made [74]. Another problem encountered with a reference redox system is caused by the fact that often the redox system is added to the electrolyte under investigation in the presence of the substrate (e.g.,  addition of fc to the solution at the end of a cyclic voltammetric experiment, recording of additional voltammograms, and determination of the fc redox potential from the resulting curves). It has recently been warned that such a procedure might introduce artifacts if redox cross-reactions occur [80]. Either the internal reference redox system must be chosen such that cross-reactions 0 are avoided  [80] or careful external determination of Eref has to be performed and checked for reproducibility. The use of fc as a reference redox system has also been extended to ionic liquids [81]. Apart from the redox couples mentioned earlier, decamethylferrocene [82] and cobaltocene [81c,d,83] have been suggested as reference redox systems. c. (Haber) Luggin Capillary The position of the Luggin capillary has a decisive influence on the measured potential. It includes an unavoidable resistive voltage drop ΔU within the cell electrolyte (see Figure 7.4(d)), dependent on the cell current i and the distance from the electrode surface. If ΔU represents a significant part of the measured potential, it becomes difficult to evaluate the real electrode potential. Also the application for controlling the potentiostat is complicated. Only in case of very low current densities—for example, in cyclic voltammetry at ultramicroelectrodes—ΔU may be negligible. In preparative electrolysis, ΔU can be estimated, for example, by calculations using the conductivity of the electrolyte or by measurements at different distances from the electrode surface. Then, a compensation in the electrical circuit of the potentiostat is possible (see manufacturer’s manuals for individual procedures). Furthermore, the fraction of all resistive voltage drops within the electrode potential is detectable by fast current interruption measurements (shutdown within about 1 µs). Any resistive voltage drop disappears immediately, while all other potentials and overpotentials stay at least some microseconds nearly unchanged. Comfortable potentiostats can correct the electrode potential automatically based on this method. Another method, compensation by positive feedback techniques [84], is common in commercial instruments since many years.

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The discussed interdependence between current density and electrode potential implicates vice versa that an exact measurement is possible only at a location in front of the working electrode with a well-defined current density. However, usually the Luggin capillary itself disturbs the local current density. A classical standard for the Luggin capillary dimension and position with minimal influence on the local current density is elucidated in Figure 7.4(e) (d ≈ 0.5–1 mm) [3]. It may be difficult to adjust the distance. A minimum distance like that in Figure 7.4(f) causes inaccurate results. The design in Figure 7.4(g) is helpful for a well-defined position and to prevent that gas bubbles from gas evolution at the electrode can enter the orifice. A blocking of the Luggin capillary by gas bubbles will be detrimental because in this case the control of the potentiostat is interrupted and its voltage and current limits may be reached, probably causing a damage in the cell. The position in Figure 7.4(h) is possible if a borehole can be applied into the electrode (however, also here the local current density is changed). A nearly optimal position demonstrates Figure 7.4(i), if it is correctly realized as, for example, in the FlexCell® of Gaskatel GmbH [85]. Here, the Luggin capillary is a fine borehole in the cell wall and ends in the cell electrolyte close to the electrode surface. No perturbation of the current density distribution is caused (a complete isolation from the electrode is necessary for reliable results). A Luggin capillary position behind the working electrode as in Figure 7.4(e) is incorrect because the potential is related to a too small current density (also in case of perforated or grid electrodes, the current density on the backside is smaller than that in the front of the electrode). The potential measurement needs a high-impedance input (measurement current typically 60), those with medium dielectric constant (20 < D < 50), and those with low dielectric constant (D < 13). In the first group are solvents like water, formamide, N-methylamides, and PC; the second one comprises compounds like acetonitrile, dimethylformamide (DMF), dimethyl sulfoxide (DMSO), methanol, nitromethane, and ammonia; whereas in the last group belong solvents like acetic acid, ethylenediamine (EDA), methylamine, tetrahydrofuran (THF), dioxane, and methylene chloride. The solvents in the last group require a higher concentration of supporting electrolyte in order to acquire a reasonable conductivity. For acetonitrile, propylene carbonate, 4-butyrolactone, DMSO, DMF, and similar amides, the maximum conductance for fully ionized salts occur at approximately 1 M concentration and is about 10 −2 Ω−1 cm−1, comparable to an aqueous salt solution of 0.5–1 M. The rate constant for the heterogeneous electron transfer to aromatic compounds is usually somewhat higher in MeCN than in DMF, which indicates that not only the dielectric but also the dynamic properties of the solvents influence the rate of radical anion formation [A70]. III.A.1.d Dissolving Power The choice of solvent is often a result of a series of compromises. Very few solvents have good dissolving power for both organic substrates and inorganic salts. The good electrochemical properties of water are marred by the lack of its ability to dissolve many organic systems. This ability may be raised by using water in a mixture with an organic solvent like ethanol, acetonitrile, DMF, or dioxane, or by using a “hydrotropic” supporting electrolyte [A71] like tetraalkylammonium p-toluenesulfonate. Tetraalkylammonium salts are soluble in most polar solvents and even in less polar liquids, such as chloroform and methylene chloride. Some polar solvents, such as acetonitrile, DMSO, and DMF, combine good dissolving power for both organic compounds and a variety of salts.

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TAbLE 7.2 Dielectric Constants and Accessible Potential Range for Some Solvents Dielectric Constant 80 33 >84 6.2 37.5 36.7 32 30 23 12.5 13 46.7 44 69 36.7 7.4 16.6 8.9

Reference Electrode

Supporting Electrolyte

Anodic Limit (V)

Pt

SCE

HCIO4

1.5

Pt Pt Pt Pt Pt

SCE Ag/Ag+ Hg pool Hg pool Ag/Ag+

NaOAc LiClO4 LiClO4 LiClO4 LiCO4

2.0 2.4 1.5 1.4 0.8

[A61]a [A62] [A63] [A63] [A62]

C Graphite Pt Pt Pt Pt Pt Pt Pt

SCE Ag/Ag+ SCE SCE SCE SCE Ag/Ag+ Ag/AgCl SCE

TEAP LiClO4 NaClO4 NaClO4 TEAP LiClO4 LiClO4 TBAPF6 TBAP

0.1 1.4 0.7 2.3 1.7 2.7 1.8 2.1 1.8

[A23] [A65] [A66] [A67] [A68] [A35] [A39] [A39] [A69]

Working Electrode

Reference Electrode

Supporting Electrolyte

Hg Hg Hg Hg Pt Hg Hg Hg Hg,Pt Hg Hg Hg Pt Hg Hg Pt Pt Pt

SCE Hg pool Hg pool SCE Ag/Ag+ Hg pool Hg pool Ag/Ag+ Hg pool SCE Hg pool SCE SCE SCE SCE Ag/Ag+ Ag/AgCl SCE

TBAP TEAB None TEAP LiClO4 TEAP TEAP LiClO4 TBAI TEAP LiClO4 TEAP NaClO4 TEAP LiClO4 LiClO4 TBAPF6 TBAP

Cathodic Limit (V) −2. −2. −0.7 −1.7 −3.5 −3.5 −3.3 −3.6 −2.3 −2.65 −1.7 −2.8 −4 −2.5 −1.2 −3.6 +0.4 −1.7

[A29] [A62] [A63] [A63] [A64] [A3] [A23] [A65] [A66] [A67] [A68] [A35] [A39] [A69]

The limit for the accessible potential for a solvent depends on many factors, such as supporting electrolyte, electrode material, and magnitude of permissible current density, and the numbers given are cited as illustrative values only. DMF, dimethylformamide; NMP, N-methylpyrrolidone; HMPA, hexamethylphosphotriamide; DMSO, dimethyl sulfoxide; sulfolane, tetramethylene sulfone; PC, propylene carbonate; THF, tetrahydrofuran; HFP, hexafluoro-2-propanol; TBAP, TBAI, tetrabutylammonium perchlorate or iodide, respectively; TEAP, TEAB, tetraethylammonium perchlorate or bromide, respectively; TBAPF6, tetrabutylammonium hexafluorophosphate. a Reference numbers.

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Organic Electrochemistry

Water Methanol H2SO4 (96–99%) CH3COOH CH3CN DMF NMP HMPA NH3 (H2NCH2)2 Pyridine DMSO Sulfolane PC CH3NO2 THF HFP CH2C12

Working Electrode

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Besides the dissolving power of a solvent toward substrate and supporting electrolyte, the ability to solvate intermediate cations and anions is also important. The chemical properties of charged species are dependent on whether there are formed tight ion pairs, solvent-separated ion pairs, or symmetrically solvated ions. Redox potentials of ions are thus dependent on solvent (see, e.g., Reference A72). A knowledge of the donor and acceptor values [A53,A54] of a solvent is thus helpful in predicting its properties as medium. Also, the autoprotolytic constant of a solvent is of interest in this respect. A combination of the good conductivity of water and the dissolving power of an organic solvent may be exploited by using an emulsion [A73–A84]. A simple emulsion is useful only if the substrate has a reasonable solubility in water and thus only a limited advantage compared with a suspension of the substrate. For reactions in which the product is much more soluble in an organic phase than the reactants, an emulsion may continuously extract the product and protect it against further reaction. Emulsions are made from an aqueous phase, an organic phase, and a surfactant; they may be “water in oil”, “oil in water”, or a bicontinuous microemulsion. The first two types have one continuous phase, whereas the bicontinuous microemulsion has both the aqueous and organic phases as continuous phases; it consists of a dynamic intertwined network of the two phases with a surfactant monolayer at the interphase [A80–A82]. Bicontinuous microemulsions are nontoxic and have a large interfacial area facilitating an intimate mixing of polar and nonpolar reactants, good solubilization of polar and nonpolar compounds, and high conductivity. Bicontinuous microemulsions could thus be useful for electrochemical synthesis as an alternative to the use of organic solvents [A80–A82]. Utilization of a mediator and a phase transfer catalyst may enhance the reaction rate considerably; in a bimolecular reaction, the rate may be enhanced by preconcentration of a mediator by adsorption to the electrode in a microemulsion [A81,A82]. When a desired bimolecular reaction involves ions (including radical ions), the nature of the surfactant (cationic, anionic, neutral) may play a role in the product distribution [A83]. Ultrasound-assisted electrolytic reduction of emulsions of activated unsaturated systems provides a method for hydrogenation of water-insoluble materials in an aqueous environment [A85]. The effect of ultrasound on electrochemical reactions in emulsions may vary depending on the reaction; in some cases, solubilization of an insoluble reaction product is furthered, whereas in other cases, the heterogeneous rate constant is influenced [A84]. When cells with solid polymer electrolytes (SPE) are used, no electrolytes need to be dissolved in the medium, and solvents, which usually cannot be employed in electrolysis, may be used [A85–A90]. III.A.1.e Temperature Range and Other Factors It is desirable that the solvent be a liquid in a convenient temperature interval and that the vapor pressure be not too high at the working temperature to avoid excessive loss of solvent. This is not an important point since the electrolysis is often carried out in a closed system anyway. The vapor pressure may also sometimes be lowered considerably by working with a high salt concentration. Too low a vapor pressure of the solvent is inconvenient, as it may be difficult to remove the solvent during the workup. The viscosity of the medium influences not only the mass transfer but also the rate constant of the heterogeneous electron transfer [A91]. A low viscosity is preferable both from the point of view of diffusion and from considerations of pumping in flow cells. For some kind of electroanalytical work, however, a high viscosity is preferable. Diffusion coefficients in some solvents useful for electrolysis have been published [A92]. The toxicity and odor of the solvent warrant consideration. Proper design of the apparatus and careful handling of the solvent make these properties less objectionable. The chirality of a solvent, S,S-(+)2,3-dimethoxy-l,4-bis(dimethylamino)-butane, has been used to induce chirality in reduction products [A93].

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III.A.1.f Purification Few, if any, organic solvents are obtained commercially pure enough for electrochemical work. However, when selecting a method for purification of a chosen solvent, one should realize that it is seldom necessary to lower the concentration of all impurities. An aprotic solvent used for a reduction should especially be purified for electrophiles, whereas nucleophiles (bases) should be removed from a solvent for oxidations. Solvents obtained from different sources may have different impurities, and even in solvents from the same source, the kind and concentration of impurities may change from batch to batch. Water is always an impurity in an organic aprotic solvent. The influence of a given water concentration depends on the stability of the complexes between water and the solvent; water is, for example, less firmly bound to MeCN than to DMF or DMSO. When one wants to remove a given impurity, there is always a danger that the reagent or method introduces other impurities or decomposition products; some solvents, such as DMF, decompose somewhat on distillation at ambient pressure. The storage of a solvent is important, and ideally, one should use the solvent immediately after purification. Some solvents are decomposed by light, and most solvents absorb some water if given the chance. Methods for purification of some aprotic solvents and tests for impurity recommended by IUPAC have been published [A94].

III.A.2

PROTIC SOLVENTS

It is practical to divide solvents into two groups, protic and aprotic solvents. Protic solvents are those that have protons bonded to heteroatoms and include acids, neutral solvents, and some bases. A review of solvents useful for electrochemistry has appeared [A69]. Electrochemical reactions in nonaqueous systems [A95] and the chemistry of nonaqueous solvents [A96] have been treated in monographs. III.A.2.a Acid Solvents The electrochemistry of a protonated organic compound differs in many cases from that of the neutral molecule, and the strength of the acid necessary for protonation of the substrate must be considered. Sulfuric acid and acetic acid, often containing some water, have generally been used. When high acidities are desired, fluorosulfonic acid may be used. III.A.2.a.i Sulfuric Acid Concentrated or slightly diluted sulfuric acid is in many respects an excellent medium for electrolysis. It dissolves many organic compounds, and it is sufficiently dissociated to make the addition of foreign electrolytes unnecessary. Sulfuric acid can protonate even very weak bases, such as aromatic ethers or some aromatic hydrocarbons; it may promote the formation of carbocations either by dehydration of hydroxyl compounds or by addition of a proton to a double bond. These properties are still more strongly developed in the “superacids”. Sulfuric acid has certain disadvantages: It cannot be removed by distillation during the workup. It is, therefore, often most practical to dilute the medium with water and extract the product. If neutralization is necessary, concentrated ammonia is preferable. Another disadvantage is that sulfuric acid may oxidize or sulfonate the substrate. Most of the electrolyses in sulfuric acid have been performed without control of the potential. In polarographic work, the mercury pool has been used for anodic and cathodic reactions: as anode, platinum, or lead dioxide has been employed and as cathode, lead, mercury, or platinum.

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A 1:1 mixture of sulfuric acid and ethanol was a very popular medium for electrolytic reductions around 1900. The mixture dissolves many organic compounds, but the leveling effect of the alcohol lowers its protonating ability considerably. III.A.2.a.ii Fluorosulfonic Acid This solvent is a very strong acid and may be used when stabilization of cationic species is required [A31,A97–A103]. It has above all been used as medium for the oxidation of alkanes. It is rather resistant toward anodic oxidation, the oxidation product being peroxydisulfuryldifluoride [A101]. Fluorosulfonic acid is not a conducting medium, and it is necessary to add a supporting electrolyte, such as NaSO3F or KSO3F, or an ionizable organic compound, such as acetic acid, which acts as a strong base and forms CH3CO+. Antimony pentafluoride acts as an acid in FSO3H. In fluorosulfonic acid containing sulfuric acid, anodic dissolution of platinum is reported to take place; in anhydrous FSO3H, an oxide layer is formed, which protects the platinum electrode [A103]. As reference electrode, the hydrogen electrode [A32], the Au(s)/Au+3 electrode [A31], and HgSO4/Hg in HSO3F [A30] have been used, but an internal reference system, such as perylene+/ perylene2+, has been recommended [A59] to avoid unknown liquid junction potentials. Perfluorosulfonic acids may be used similarly [A104]. They may be purified by distillation and treatment with H 2O2 [A105]. Trifluoromethanesulfonic acid has been used as a solvent for CV with sodium trifluoromethanesulfonate as supporting electrolyte, glassy carbon as indicator electrode, platinum as counter electrode, and a silver wire as pseudo-reference electrode. At v = 100 mV s−1, the accessible potential window was from +0.4 to +3.0 V vs. NHE (calibrated against Ru(bpy)33+/2+) [A106]. III.A.2.a.iii Hydrogen Fluoride Hydrogen fluoride has been used for electrochemical fluorination of organic compounds [A107– A110] and is a promising solvent for electrochemical studies. It has a high dielectric constant (80 at 0°C), is less viscous than water, is transparent for UV light to 165 nm, and is difficult to oxidize. It dissolves many metal fluorides, giving highly conducting solutions; it also dissolves many organic substances. The main disadvantages are its poisonous effect, its ability to attack glass, and its relatively low boiling point, 19.5°C. The low boiling point gives it a tendency to form bubbles, which must be taken into account in the construction of cells. This and other practical problems have been discussed [A111]. Hydrogen fluoride can be handled in an apparatus of suitable metals (copper, nickel, magnesium, or aluminum, which all form a protective fluoride coating, or platinum) or plastic materials [especially polypropylene, Teflon, and polyvinylidene fluoride (Viton)]; polychlorotrifluoroethylene (Kel-F) can be made into transparent windows. A capillary for a dropping mercury electrode may be made from Teflon [A112]. Hydrogen fluoride is obtained commercially in steel cylinders in a purity of 99.5%. The impurities may be removed by distillation [A113] or electrolysis [A114]. During the electrolytic removal of water, the explosive F 2O is formed, which must be taken into consideration [A110]. The useful potential range in the cathodic direction is rather limited owing to the evolution of hydrogen, whereas the anodic limit is about 2.6 V (SCE). Sodium or potassium fluoride as supporting electrolyte may be used; the addition of these salts lowers the protonating power of the solvent. The solubility of NaF is about 30 g NaF per 100 g HF; other fluorides, such as LiF, KF, CsF, NH4F, AgF, and TlF, are very soluble in liquid HF. The acidity of HF may be enhanced by the addition of BF3 or SbF5. The addition of amines to hydrogen fluoride makes the fluoride ion a better nucleophile and reduces the acidity of HF; Et3N/3HF and pyridine/HF are commercially available. Et3N/3HF may be used as such or mixed with MeCN as a medium for electrochemical fluorination (Chapter 20) [A115,A116].

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Several types of reference electrodes may be used. The hydrogen electrode has been shown to behave reversibly in HF [A44,A45] and has been employed as a reference electrode both in electroanalytical and in preparative work [A45], and mercury fluoride [A41,A117] and copper fluoride [A43,A111] have also been used as such. A comparison of these reference electrodes pointed to the hydrogen electrode (H2/Pd) as the most convenient [A118]. III.A.2.a.iv Trifluoroacetic Acid Trifluoroacetic acid (TFA) has mainly been used as a medium for oxidations [A119]. It has a dielectric constant of 8.4 [A120,A121] and a boiling point at 72.5°C; although it has a low conductivity, it is possible to conduct electrolysis in it without using a supporting electrolyte [A122]. Its main advantage is that many cation radicals show considerable stability in this medium. Sometimes it is used in a mixture with its anhydride and methylene chloride [A49]. III.A.2.a.v Acetic Acid Acetic acid is a good solvent for many organic compounds and some inorganic salts. It has a rather low dielectric constant, which makes it necessary to have a considerable electrolyte concentration (≥0.5 M) to obtain a reasonable conductance. Acetic acid may be considered as solvent for acetoxylations [A61,A123,A124] and for reactions that require an acidic medium other than sulfuric acid. The acidity may be enhanced by adding perchloric acid to acetic acid containing a suitable amount of acetic anhydride, which reacts with the introduced water. Usable supporting electrolytes include NaOAc, NH4OAc, LiCl, HCl, H2SO4, HClO4, NaClO4, Bu4NClO4, and Bu4NBF4. The choice of supporting electrolyte may influence the product distribution. Thus, anodic acetoxylation gives different results in NaOAc/HOAc and in Bu4NBF4/HOAc. Several types of reference electrodes work well in acetic acid. The chloranil electrode [A26] behaves reversibly in this medium and so does the analog of SCE [A65] [Hg/Hg2Cl2(s), LiCl(s), HOAc]; the mercury pool [A125] and the Ag/AgCl electrode [A24] can also be used. The cathodic limit is about −1.7 V (SCE), and the limiting reaction is evolution of hydrogen. The anodic range ends at about 2.0 V (SCE) when NaOAc is the electrolyte [A61]. The limiting reaction is probably the discharge of acetate ions. Acetic acid may be purified by partial freezing [A61] or by distilling it from CrO3 [A65,A125] in the presence of acetic anhydride. CrO3 oxidizes impurities and acts as an acid catalyst for the reaction between water and acetic anhydride. An excess of acetic anhydride is necessary to force the equilibrium 2AcOH ⇌ Ac2O + H2O sufficiently to the left [A126]. Modifications of the method have been described [A23,A125]. III.A.2.a.vi Other Acids Methanesulfonic acid [A127] and formic acid have also been used as solvent for electrolysis; superacids offer possibilities for unusual electrode reactions. III.A.2.b Neutral Solvents This group comprises water, mono- and polyvalent alcohols, and monoethers of polyvalent alcohols. Often mixtures of water and alcohols are used. Such mixed solvents retain the favorable characteristics of water to some degree but raise its ability to dissolve organic substrates. III.A.2.b.i Water Water is the solvent of choice for electrolysis, unless some of its properties are undesired or inadequate. An aprotic solvent is sometimes preferable to water since a reaction mechanism of an electrode reaction in an aprotic, organic solvent is often simpler to elucidate than in water. Adsorption phenomena are often a complicating factor in aqueous solution, but much less so in many organic solvents. The usable potential range of water in the anodic direction is quite

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limited, and the most important factor, the solubility of many organic compounds in water, is quite unsatisfactory. The use of “hydrotropic” salts (Appendix B, Section III.B.1) may enhance the solubility of organic compounds in an aqueous medium. Another way of combining the good conductivity of aqueous solutions with the better dissolving power for organic substrates is the use of emulsions [A73–A84], preferentially with a mediator and a phase transfer reagent. III.A.2.b.ii Methanol The electrochemical behavior of methanol is rather similar to that of water. Methanol is well suited for anodic reactions of the Kolbe type and for methoxylations; for reductions, it may be used when the solubility of the substrate in water is too low. The dielectric constant of methanol is fairly high (ε = 33); its liquid range (−98 to 64°C) is convenient, and it is easy to remove during the workup. Quite a few electrolytes are soluble in methanol, such as NH4Cl, LiCl, HCl, KOH, KOMe, NaClO4, and tetraalkylammonium salts. For reductions, a solution of HCl in MeOH is convenient as it has a high conductance and both components are easy to get rid of during the isolation of the product. Reference electrodes that can be used in water work in most cases also in MeOH; thus the Hg/Hg(I) and Ag/Ag(I) electrodes may be used. The useful potential range for large-scale electrolysis is about the same as for water, but methanol is less satisfactory for electroanalytical work. Methanol can in general be used as received; it can be dried by treatment with magnesium [A128]. It has been reported that methanol oxidized at the anode may diffuse to the catholyte and react with nucleophiles or EGB [A129]. III.A.2.b.iii Other Alcohols Ethanol and the higher alcohols have lower dielectric constants and poorer ability to dissolve electrolytes. They have been used mostly in mixture with water or sulfuric acid. 2,2,2-Trifluoroethanol is more acidic than ethanol and can be used when strong bases are undesired [A128]. III.A.2.c 1,1,1,3,3,3-Hexafluoropropan-2-ol 1,1,1,3,3,3-Hexafluoropropan-2-ol (HFP) is a solvent [A39] with unusually low nucleophilicity; it has a high ionizing power, a high hydrogen bonding strength, and a low hydrogen bonding acceptor strength, which makes it a good solvent for investigations of cation radicals. It may be useful for electropolymerization. It dissolves many polar compounds, but is less good in dissolving nonpolar substrates. It has a liquid range from −5 to 58.6°C, a dielectricity constant of 16.7, pKA in water 9.30, and pK A in DMSO 18.2. It is rather volatile and should be handled in a well-ventilated hood. Tetrabutylammonium hexafluorophosphate or tetrafluoroborate has a sufficient solubility in HFP to give a good conductivity. Aqueous Ag/AgCl or SCE has been used in connection with a suitable low-leaking salt bridge. The reversible potential ferrocenium/ferrocene couple in HFP was found [A39] to be 0.05 V vs. Ag/AgCl, to be compared with that in CH2C12, 0.43 V vs. Ag/AgCl. The potentials in HFP generally differ appreciably from those in, for example, MeCN. HFP of highest quality may be used as received; it is rather expensive as a solvent for electrochemistry and regeneration of the solvent could be attractive, but as HF might be formed during regeneration, the procedure should be carefully considered. However, in the author’s laboratory, solvents from CV (HFP/TBABF4), depending on the substrates, have been regenerated by distillation at 30°C/½ atm after use with a loss of 10–15%. III.A.2.d basic Solvents Basic, protic solvents include ammonia and primary and secondary amines. These solvents are primarily of interest for an organic electrochemist because of their ability to solvate electrons, and solvated electrons have special reducing properties (Chapter 29).* They also permit reductions in a protic medium in the presence of a very strong base, the conjugate base of the solvent. * This chapter reference refers to the fourth edition.

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III.A.2.d.i Ammonia The use of ammonia as a solvent for electrochemical reactions presents some problems due to its inconvenient liquid range (−77.7 to −33.4°C); this difficulty can be coped with in different ways. The electrolysis may be made in a suitable low-temperature thermostat [A131], the cell can be equipped with an efficient reflux condenser [A132], a high-pressure apparatus can be used [A133] (ammonia has even been used in its supercritical phase as a medium for electrochemical investigations [A134]), or strong solutions (5–10 M) of various salts, such as NH4SCN, NH4I, LiNO3, LiClO4, NaSCN, and NaI, can be used [A135] as they form a medium in which the vapor pressure of ammonia at ambient temperature is below 1 atm. The easy purification and low price of ammonia compared with the alternative amines favor the choice of ammonia for large-scale preparations. When using a divided cell, a tube must connect the anode and cathode compartments for equilibration of the pressure. Ammonia has a medium–high dielectric constant (23.7 at −36°C) and good dissolving power toward inorganic salts but less for nonpolar organic compounds. It can act as both an acid and a base; the strongest possible acid in ammonia is ammonium ion and the strongest base the amide ion. The pK values of some weak acids have been determined in NH3 at −60°C [A136]. The combination of low proton availability together with the low temperature usually employed makes ammonia a suitable medium for electrochemical investigations of strongly basic species, such as dianions [A137,A138]. The low tendency to donate hydrogen atoms, due to the strength of an N–H bond, is important for many investigations [A139–A141]. Besides the electrolytes mentioned earlier, such salts as NH4Cl, NH4NO3, KNO3, and NaClO 4 may be used as supporting electrolyte; the tetraalkylammonium salts are generally rather sparingly soluble in ammonia. Most electrolyses in ammonia, except in voltammetric studies, have been performed without a reference electrode; as such, a Zn(Hg)/ZnCl2 [A13] or Pb/Pb2+ electrode [A3,A131,A142] may be used. In polarography, the mercury pool electrode has been applied [A3]. Ammonia has been employed mostly for cathodic reactions, but some oxidations [A132,A135] have been carried out in this medium, although the potential range in the anodic direction is quite small. The anodically limiting reaction is oxidation to nitrogen and protons [A143]; the cathodically limiting reaction is the transfer of electrons to the solvent, which occurs at about −2.3 V (vs. Hg pool electrode) in a saturated solution of TBAI. In the electrolytic generation of solvated electrons, the potential is determined by the surface concentration of electrons and no external reference electrode is needed. Ammonia is purified by distilling it after treatment with sodium; the distillation from sodium may have to be repeated. The threshold limit value (TLV) is 25 ppm. III.A.2.d.ii Methylamine Methylamine behaves electrochemically much like ammonia but has a somewhat more convenient liquid range (−93.45 to −6.3°C), although its odor and low boiling point require special considerations. Its dielectric constant is rather low (11.4 at −10°C), but sufficient conductivity for large-scale electrolysis can be obtained using LiCl as supporting electrolyte. Like ammonia, it is both an acid and a base. Methylamine has mainly been used in large-scale electrolysis [A144,A145], and no reference electrodes have been employed so far. Reference electrodes such as the Zn(Hg)/Zn2+ electrode, which work both in ammonia and in EDA, would be assumed to be applicable to methylamine solutions. Methylamine has been employed only for reductions, especially as solvent for electrolytic generation of solvated electrons. It would be of very limited use for anodic reactions. It may be purified by distillation from sodium. Its TLV is 10 ppm.

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III.A.2.d.iii Ethylenediamine Ethylenediamine (EDA) is a primary amine with a more convenient liquid range (11–117°C) than methylamine. Its dielectric constant (ε = 12) is about the same as that of methylamine. It dissolves many inorganic salts and is a better solvent than ammonia for many organic substrates. LiCl, NaNO3, and tetraalkylammonium salts can be used as supporting electrolytes. For the electrolytic generation of solvated electrons, mainly LiCl has been employed [A146,A147]. A reversible reference electrode in EDA is the Zn(Hg)/Zn2+ electrode [A148], but the Hg pool [A15] or the aqueous calomel electrode, fitted with a suitable salt bridge, is also applicable. EDA is oxidized rather easily anodically but is resistant to reduction; the cathodic limiting current is the discharge of cations or the donation of electrons to the solvent. EDA may be purified by repeated refluxing and distillation from sodium [A149]. It is extremely hygroscopic. The TLV is 10 ppm. Other primary or secondary, aliphatic or cyclic amines may be used with similar results.

III.A.3

APROTIC SOLVENTS

The use of aprotic solvents is of interest in many cases in which aqueous solutions lack desired or have undesired properties. The scarcity of protons makes it generally simpler to elucidate the mechanism of an electrode reaction for several reasons. The intermediates (e.g., radical anions) are more stable in the absence of protons, and the reaction scheme becomes simpler. The scarcity of protons in aprotic solvents also makes it possible for added reagents (e.g., electrophiles) to compete successfully with protons for the intermediates. “Dry” aprotic solvents usually contain some water, which is especially of importance in voltammetric studies in which the concentrations of water and substrate are about equal. The influence of a certain molar concentration of water depends on the solvent, the supporting electrolyte, and the substrate. In certain solvents, such as DMF and DMSO, water is a rather poor proton donor [A150], and other impurities may be responsible for the protonation of the basic intermediates (radical anions, anions, and dianions). During preparative experiments, the impurities may be reprotonated by water or, in case tetraalkylammonium salts (except tetramethylammonium salts), are used as supporting electrolyte, by attack on the cations (Hofmann elimination). Treatment of the medium with active alumina may lower the concentration of such protonating impurities [A51,A151]. When choosing a drying agent, one must take into account that its affinity to water is different in vacuum and in a solvent, so the activity of a drying agent may be different in various solvents. When deuterium oxide is added to an aprotic solvent containing tetraalkylammonium salts to promote incorporation of deuterium in the product, the base-promoted Hofmann elimination may become a problem, as the deuterium oxide becomes diluted with water. This can largely be avoided by running the reduction in the presence of a suspension of dry silica, which prevents the medium from being basic enough for the elimination to occur; the deuterium oxide adsorbed on the silica is in most cases sufficiently acidic to prevent Hofmann elimination. However, when the parasitic reaction is hydrogen abstraction from the solvent or supporting electrolyte, silica has no effect. Controlled addition of a suitable proton donor or electrophile (reductions) or nucleophile (oxidations) is often useful in determining a reaction mechanism. The strength of a proton donor may vary from perchloric acid through acetic acid and a phenol to an alcohol; C acids, such as malonic ester, or N acids, such as urea, may also be used. Used as bases may be pyridine, carboxylate ions, alkoxides, or salts of malonic ester. Sometimes it is of interest to determine whether it is the basic or the nucleophilic properties of the compound that are important. The use of two bases with approximately the same pK values but widely differing in nucleophilicity, such as pyridine and a 2,6-dialkylpyridine, might answer the question.

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In the absence of proton donors, the “aprotic” solvent may act as a proton donor; the conjugate base of the solvent may act as a nucleophile toward substrate or intermediates. Electrochemical phenomena may also be simplified in organic solvents, compared with aqueous solutions, because of the less important role adsorption plays in organic solvents. III.A.3.a Acetonitrile Acetonitrile (MeCN) is one of the most used polar aprotic solvents for both anodic and cathodic reactions. It is an excellent solvent for many organic substrates and quite a few organic and inorganic salts; it is miscible with water. Salt solutions show a reasonably high conductivity due to its rather high dielectric constant (ε = 37). MeCN is somewhat toxic, with a recommended maximum air concentration of 20 ppm. Besides having good electrochemical properties, acetonitrile is an unusually good solvent for UV spectroscopy, being transparent until about 190 nm. It is furthermore a convenient solvent for electron spin resonance spectroscopy, which is valuable for the investigation of radicals produced during electrolysis. For use in preparative electrochemistry, its liquid range (−45 to 82°C) is sufficient for most purposes and its relatively low boiling point makes the removal of the solvent easy. On gas–liquid chromatography, tailing may occur, unless a suitable polar column is used. Acetonitrile has a wide, usable potential range in both the anodic and cathodic directions. The limit is set by the electrode reaction of the supporting electrolyte in both directions. In the presence of sodium or lithium ions, their discharge sets the limit; the sodium subsequently reacts with acetonitrile, whereas lithium does not. Quaternary ammonium ions are reduced with dealkylation. Perchlorate ions, a widely used supporting electrolyte, are discharged to perchlorate radicals, which probably rapidly form oxygen and chlorodioxide radicals [A50]. Transfer of protons in acetonitrile is a rather slow process. Its autoprotolysis constant is 33.2. In DMSO, MeCN has pK A = 31.3 [A152]. MeCN is less apt to donate hydrogen atoms than DMF. Anhydrous perchloric acid in MeCN has been prepared by oxidation of H2 to H+ at a platinized platinum electrode in a solution of TBA+ClO4− in MeCN [A153]. Useful supporting electrolytes in acetonitrile are NaClO4, LiClO4, tetrabutylammonium salts, as well as other tetraalkylammonium salts. AgNO3 and AgClO4 are very soluble in acetonitrile; AgCl is somewhat soluble, especially in the presence of an excess of chloride ions. It is possible to make CV in MeCN with TEACF3SO3 or CF3SO3Na as supporting electrolyte near the supercritical conditions, but the resistance is too high for preparative reactions [A154]. The Ag/Ag+ electrode [A155,A156], which is reversible in acetonitrile, may be used as reference electrode. Quite often, an aqueous calomel electrode with a suitable liquid junction is used [A3]; a comparison of different reference electrodes has been made [A4]. The Hg/Hg(I) is not stable in acetonitrile, as Hg(I) decomposes [A157]. A platinum electrode coated with a film of poly(vinylferrocene) is a stable reference electrode in MeCN [A158]. Commercial acetonitrile usually contains impurities, such as acrylonitrile, acetic acid, aldehydes, amines, and water. Several methods of purification [A159–A161] have been proposed: one is to distill it several times from P2O5 followed by fractionation from K2CO3 [A162]; a pure product may be obtained, but the procedure is time consuming and wasteful due to polymerization of MeCN; polymerization may be avoided by using B2O3 instead of P2O5 [A163]. CaH2 treatment of MeCN has been recommended [A141], but it does not remove traces of aromatic hydrocarbons and is a surprisingly ineffective drying agent in MeCN [A163]. For removal of acrylonitrile, which boils 4°C lower than MeCN, treatment with NaH [A164] or azeotropic distillation with ethanol [A165] has been recommended. Removal of impurities by reaction in turn with benzoyl chloride and KMnO4 followed by careful distillation gives a product suitable for both electrochemical and spectroscopic work [A166]. Alternatively, acetonitrile is treated with N2O4, then with CaH2, fractionated under N2 from CaH2, and the product further purified by letting it run through a column of activated alumina [A51,A165,A167].

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Water may be removed from commercial MeCN by azeotropic distillation [A62,A168] or by treatment with molecular sieves 3A; 3A is much more efficient than 4A for drying MeCN [A163]. The procedures recommended by IUPAC for purification are described in Reference A169. As is the case for other relatively inert solvents, it is unlikely that an all-purpose purification procedure can be devised that will give optimum results in all possible applications. Consequently, purification should be tailored to the intended use of the solvent. Pure acetonitrile is not particularly hygroscopic. The product obtained by these procedures can be stored for several months without deterioration; the water content is approximately 1 mM. It is advisable to use the same glass apparatus repeatedly for the purification and storing of acetonitrile. Electrolysis in MeCN, in the absence of added proton donor may produce the anion of MeCN (CH2CN)−, which may act as a nucleophile toward electrophilic centers [A170]. At a Pt anode, the solvent system MeCN/TBABF4 breaks down to an adduct, CH3CNBF3, whereas the cathodic breakdown mainly gives the anion of 3-amino crotonitrile; in the presence of water, the anion of acetamide is also formed and the breakdown commences at much less negative potentials [A171]. Other nitriles, aliphatic and aromatic, have been investigated as solvents for electrolysis. Butyronitrile has been suggested as the solvent of choice for low-temperature electrochemistry [A172], and propionitrile has in some cases been employed, as it is less hygroscopic than MeCN [A173]. Benzonitrile might be considered when hydrogen atom abstraction from the solvent is undesirable. III.A.3.b Dimethylformamide Dimethylformamide (DMF) is a polar solvent with a dielectric constant (ε = 37) comparable to that of acetonitrile. It is a good solvent for most organic compounds and also for many organic and inorganic perchlorates and organic fluoroborates. Owing to these properties, DMF has been widely used in electrochemical work. DMF is, however, susceptible to hydrolysis, which yields the easily oxidizable dimethylamine and the reducing agent formic acid. DMF has an acceptable liquid range (−61 to 153°C), although the boiling point is near the upper limit for its convenient removal by distillation during the workup. DMF is inferior to acetonitrile as a solvent for UV spectroscopy. Contact with the skin and vapor concentrations over 10 ppm should be avoided. For reductions, DMF has a usable potential range comparable to that of acetonitrile, but it is inferior for oxidations. In the presence of inorganic ions, their discharge is the cathodic limiting reaction, whereas it is more uncertain whether it is the solvent or the cation that is reduced in solutions of tetraalkylammonium ions. The anodic limiting reaction at a Pt electrode is an oxidation of DMF, which involves the removal of an electron from the amide nitrogen. Its autoprotolysis constant is 29.4. DMF is a better hydrogen atom donor than MeCN and DMSO. As supporting electrolyte, tetraalkylammonium perchlorates or fluoroborates are mostly used, but LiCl or NaClO4 may also be employed. The reference electrode may be a modified Ag/AgCl electrode [A10,A12], a Cd(Hg)/Cd2+ electrode [A10,A12], Na(Hg)/Na+ electrode [A11], or silver/ silver cryptate electrode [A174]. For many purposes, the commercial reagent-grade DMF, dried with molecular sieves, can be used; DMF decomposes on distillation at atmospheric pressure. The most convenient way to purify DMF is probably to let it pass through a column [A165] of alumina (e.g., ICN Alumina N-Super I). Other methods involve drying (CuSO4 [A167], which also removes amines), azeotropic distillation with benzene, or percolation through molecular sieves [A176] followed by fractional distillation at a reduced pressure. DMF is difficult to obtain in an anhydrous form, but for most purposes, a small water content is not critical. An impurity of N-methylformamide may act as a proton donor or otherwise interfere in the follow-up reactions [A177]. DMF should be stored in the dark under nitrogen. The different purification methods have been compared and discussed [A178].

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III.A.3.c N-Methylpyrrolidone DMF has certain disadvantages that can be traced to its hygroscopic properties and rather easy hydrolysis; N-methylpyrrolidone (NMP), which is an N-disubstituted amide just like DMF, is less apt to hydrolyze owing to its cyclic structure. Furthermore, the product of hydrolysis is not a reducing agent as is formic acid formed from DMF. Although only a limited number of experiments in NMP have been done [A63,A179], it seems probable that NMP will replace DMF in a number of cases because of its higher stability [A180]. Vapor concentrations over 100 ppm should be avoided. Like DMF, NMP is of limited value for oxidations but useful for cathodic reactions. Its dielectric constant is 32, and it is able to dissolve the same types of salts and organic compounds as DMF. NMP has a liquid range of from −24 to 205°C; the useful potential range is approximately the same as for DMF. Mercury in contact with its oxidation products has been used as reference electrode; it is more stable than the Ag/Ag+ electrode [A63,A179]. Other organic amides, such as formamide, N-methylformamide, N-methylacetamide, and tetramethylurea (TMU), have been investigated as solvents for electrochemical use primarily due to their high dielectric constant; none of them, however, seem to offer distinct advantages over DMF or NMP, and in some respects, they are inferior. N,N-Dimethylacetamide may be purified by distillation in vacuum [A181]. III.A.3.d 3-Methyl-2-Oxazolidinone (3M2O) 3M2O is a colorless, nonvolatile (kp, 87–90°C) solvent with a high dielectric constant (77.5). It may be purified by vacuum distillation over BaO. With tetrabutylammonium perchlorate (TBAP) as electrolyte, the cathodic limit is −2.9 V (aqueous SCE); the anodic limit at a Pt anode is +1.3 V. The suitability of 3M2O for voltammetry has been investigated [A180]. III.A.3.e Hexamethylphosphoramide (HMPA) HMPA is a polar solvent that is remarkable for its ability to solvate electrons. Even in a mixture with ethanol (67 mol% ethanol and 33 mol% HMPA), reductions via solvated electrons may take place [A180–A184]. It has been suggested that HMPA is selectively adsorbed at the electrode, thus providing a layer of low proton activity near the electrode. The liquid range of HMPA is 7.2–235°C; its dielectric constant is 30 (20°C). HMPA solvates cations strongly, whereas anions are less solvated. Unlike carboxamides, such as DMF, it is not attacked by aqueous alkali at t < 80°C, and it is very resistant toward nucleophilic attack [A184] but forms peroxides under the influence of light and oxygen [A183]. HMPA is completely miscible with water; it can be extracted from water by chloroform, which forms a complex with HMPA [A185], K(CHCl3/H2O) = 5.5 [A186]; HMPA is somewhat more basic than DMF. HMPA dissolves such salts as LiCl, LiClO4, NaClO4, and R4NClO4; the Ag/Ag+ may be used as reference electrode [A64]. The solvation effect of the cation of the supporting electrolyte in HMPA has been investigated [A187]. Using R4NClO4, the accessible potential range is from +0.14 to −3.35 V (Ag/Ag+) at a mercury electrode [A183]. HMPA can be purified by fractionation under vacuum, boiling point 88–92°C (3 mm), 68–70°C (1 mm). Reflux over BaO, followed by fractionation under vacuum, reflux over sodium, and finally vacuum distillation have been recommended for purification [A187]. Alternatively, fractional distillation over CaH2, followed by purification through a column of activated alumina and standing over pyrene and sodium, produced a product that after distillation on a vacuum line contained only a negligible amount of impurities [A188]. The purification of HMPA has been critically discussed [A189]. Just like DMF, HMPA is not suitable for anodic reactions, since the amide nitrogen loses an electron easily. For cathodic reactions, the useful potential range extends, in the presence of LiCl, to the value at which electrons are given off to the solution. HMPA has not been used widely in electrochemical experiments, and although its special properties are useful in electrochemistry, it

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should be used only when sufficient precautions have been taken with regard to its carcinogenic and other toxic properties [A66]. Tetramethylurea (TMU) may in some cases be used as a nontoxic replacement for HMPA [A190]. TMU may be purified by distillation from CaH2. III.A.3.f Pyridine Pyridine is a reasonably strong base and a good nucleophile. Although its dielectric constant is rather low (ε = 12), many salts are soluble in pyridine and the solutions have good conductivity. In contradistinction to most bases, pyridine, which is a π-electron-deficient heteroaromatic compound, is rather resistant toward oxidation. Pyridine has a liquid range from −41 to 115°C; it is miscible with water and a good solvent for many organic compounds. Such salts as LiNO3, LiCl, LiClO4, NaI, KSCN, and tetraalkylammonium salts are soluble in pyridine. As a reference electrode, the Ag/Ag+ [A17] or the Hg pool electrode can be employed. Solute–solvent interactions have been discussed [A65]. The useful potential range is from about −2.3 V [A17] (versus Ag/Ag+ electrode, DME, Bu4NI) to  +1.4 V vs. Ag/Ag+, graphite electrode/LiClO4 [A65]. The cathodic limiting reaction might be the discharge of the cation, whereas an electron abstraction from pyridine is the anodic limiting reaction, in which probably 2-pyridylpyridinium ion is formed. Besides being used as a solvent, pyridine has been employed as a nucleophile that is oxidizable with difficulty [A156]. Reagent pyridine is satisfactory for most purposes; it may be dried with molecular sieves, type 5A. A recommended procedure for obtaining pure pyridine has been published [A191]. The TLV is 5 ppm. III.A.3.g Dimethyl Sulfoxide (DMSO) DMSO is an excellent solvent for many inorganic salts and organic compounds. It is difficult to reduce and fairly resistant to electrolytic oxidation. Its dielectric constant is high (ε = 47). It thus has many of the qualities desirable for a solvent for electrolysis, and it shows promise of being one of the most important electrochemical media [A192]. The liquid range is from 18°C to 189°C, which makes it somewhat inconvenient to get rid of DMSO in the workup. When used as solvent for electrolysis, it must be considered that DMSO is not always inert but has a fair reactivity in certain reactions. DMSO is unfit for UV spectroscopy. Its autoprotolysis constant is 31.8. DMSO has low toxicity; it is, however, transported through the skin very rapidly. Thus, minutes after applying a drop of DMSO to the palm of the hand, its sweetish taste can be detected in the mouth. It may carry with it dissolved toxic materials that otherwise would be unable to penetrate the skin. Many salts are soluble in DMSO, so the choice of supporting electrolyte is less restricted than in most other nonaqueous solvents. In general, perchlorates, even KClO4, nitrates, and halides, are soluble, whereas fluorides, cyanides, sulfates, and carbonates are not; thus, not only NaClO4, LiCl, NaNO3, and tetraalkylammonium salts can be used but also such salts as NH4PF6 and NH4SCN. The ability of DMSO to solvate ions is also of importance in the indirect electrolytic hydrodimerization of, for example, acrylonitrile using Na(Hg) [A193]. The equilibrium constant of the self-ionization of DMSO is 5 × 10 −8 (25°C); the conjugate base (CH3SOCH2)−, which is formed when DMSO acts as a proton donor during a reduction, is a rather strong base and a fairly good nucleophile, which may attack electrophilic centers or radicals [A194]. The potential range of DMSO in the cathodic direction is at an Hg electrode limited by the discharge of the cation. Potassium and sodium react with DMSO, whereas lithium does not; tetraalkylammonium ions are reduced to about −2.8 V (SCE) [A195]. The anodic limiting reaction at a Pt electrode is the oxidation of DMSO; probably one of the nonbonding electrons is lost, but the limiting reaction has not been investigated yet. As reference electrode [A10] for polarographic work, an aqueous or methanolic calomel electrode, connected to the DMSO solution with a suitable salt bridge to avoid contamination, has been

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used. The most stable reference electrodes in DMSO seem to be the amalgam electrodes, such as Tl(Hg)/TlCl [A21] or Li(Hg)/LiCl [A196]. DMSO is somewhat hygroscopic, and the main impurity in a good commercial grade of DMSO is water; the odor is due to dimethyl sulfide. For many purposes, DMSO can be used as received or after drying either by shaking with CaO–MgO or by passing through a column of molecular sieves. For further purification, fractional distillation or crystallization may be used. It has been recommended [A197] that a fractional crystallization be done followed by two fractional distillations in vacuo, one through a column packed with activated carbon and one through a spinning band column. The pure DMSO must be handled without contact with humidity. The different purification procedures have been discussed [A198]. III.A.3.h Sulfolane Among sulfones, sulfolane (tetrahydrothiophen-1,1-dioxide, tetramethylene sulfone) shows promise for electrochemical use. It has a wide liquid range (28.6–285°C); on dissolving 0.1 M salts in sulfolane, its melting point is depressed below 25°C. Sulfolane has a relatively high dielectric constant (43 at 20°C), it is chemically stable toward both reducing and oxidizing agents, and it has a fair dissolving power for inorganic salts. One of its drawbacks is that it is rather hygroscopic. The accessible potential range is large in both the anodic and cathodic directions. The limiting cathodic reaction seems to be the discharge of the supporting electrolyte, whereas the limiting anodic reaction probably is the oxidation of the solvent. With NaClO 4 as supporting electrolyte, the anodic limit at a Pt electrode is about 2.3 V and the cathodic limit about −4 (versus Ag/Ag+ electrode) [A67]. As a supporting electrolyte, NaClO4, LiClO4, NH4PF6, and quaternary ammonium compounds can be used. The Ag/Ag+ [A67,A199] or Ag/AgCl [A200] electrode or an aqueous SCE [A201] with a suitable salt bridge can be used as the reference electrode. Sulfolane may be purified by fractional crystallization followed by fractional distillations at reduced pressure [A201]; alternatively, treatment with aqueous hydrogen peroxide, followed by extraction with CH2C12 and fractional distillation, may be employed [A67]. The purification procedures have been discussed [A202]. Sulfolane is an alternative to acetonitrile or nitromethane as a solvent for anodic reactions; in reductions, it generally offers no advantages over such solvents as MeCN, DMF, or DMSO, but for anodic fluorination, sulfolane might be better than MeCN, as a competing acetamidation is avoided (Chapter 20). Among other sulfones, dimethyl sulfone has been used as electrochemical medium [A203,A204]; its high melting point, 109°C, makes it less suitable for organic compounds. III.A.3.i Propylene glycol Sulfite A series of other sulfur-containing solvents has been investigated for electrolytic use. Of these, propylene glycol sulfite [A205], a derivative of 1,3,2-dioxthiol, was found most suitable. It has a dielectric constant of 33, forms solutions with good conductivity with such salts as LiClO4, and does not react with lithium. It has not yet been used for organic electrolysis. III.A.3.j Nitromethane Nitromethane is one of the few solvents useful for anodic reactions, among them anodic coupling of aromatic hydrocarbons. Its application for cathodic reactions is limited, but in some cases in which its unreactivity toward certain active halogen compounds is valuable, it may be used. The liquid range of nitromethane is −29 to 101°C; its dielectric constant is 37, but the dissociation of salts is not as high as could be expected. Nitromethane dissolves rather few inorganic salts; LiClO4 [A206], Mg(ClO4)2, and tetrabutylammonium salts can be used as supporting electrolyte. The Ag/Ag+ or Ag/AgCl electrode can be employed as the reference electrode [A35].

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At a platinum electrode, potentials of about 3 V (vs. Ag/AgCl) can be reached, and the limiting reaction is the discharge of the supporting electrolyte. The useful range in cathodic direction is very dependent on the water content. Nitromethane has been purified by repeated washings with bicarbonate, sulfuric acid, and water, followed by drying and fractional distillation; the water content of the product is 5–8 mM. Alternatively, fractional distillation at reduced pressure followed by repeated fractional freezing may be employed [A207]. The TLV is 100 ppm. III.A.3.k Nitrobenzene Nitrobenzene has been used for electrolysis [A208]; it was found that certain radicals were rather stable in this solvent. Nitrobenzene has a liquid range from 5.7 to 210.9°C; Bu4NClO4 may be used as supporting electrolyte. An aqueous SCE separated from the solution by a suitable bridge with porous glass has been used as reference electrode. Nitrobenzene may be purified by passing it through a column of alumina followed by a distillation in vacuo. III.A.3.l Propylene Carbonate (PC) PC is of potential use in organic electrochemistry [A68,A209,A210]. It has a high dielectric constant (ε = 69) and a wide liquid range (−49 to 242°C), but it is more reactive than acetonitrile, and the potential limit for anodic reactions is lower than for acetonitrile. It has so far been used mostly for electrochemical generation of cation radicals [A68]. The cathodic limiting reaction of PC containing LiClO4 has been shown [A211–A214] to be a reduction to propene and carbonate. This reductive elimination is analogous to the reduction of vicinal halogen compounds to alkenes. The anodic reaction is an oxidation to CO2 [A213,A214]. PC decomposes slowly on standing to allyl alcohol and CO2; after standing for a year, small amounts of allyl alcohol could be detected. Protons increase the rate of decomposition [A215]. With LiAsF6, the oxidation limit of PC is 4.0 V vs. Li+/Li [A216]. As a reference electrode, an Li/LiCl electrode, which in the presence of excess tetrabutylammonium chloride is an electrode of the second kind, may be used [A217]. An Ag/AgClO4 reference electrode with an NaClO4 salt bridge has been used [A33]; AgCl/Ag [A218] and TlX/Tl [A219] reference electrodes have also been employed. The purification of PC has been discussed [A220]. III.A.3.m benzene and Chlorobenzene It has been possible to use these nonpolar solvents for electroanalytical measurements by employing a minute working electrode [A221], an ultramicroelectrode, or using a benzene solution containing a crown ether (15-crown-5) and sodium tetraphenylborate [A222]. Tetrahexylammonium perchlorate (THAB) and tetrabutylammonium tetrafluoroborate have been used as supporting electrolyte in benzene and chlorobenzene, respectively. The latter was suggested [A221] to be an excellent solvent for the study of reversible oxidations and reductions of aromatic compounds. The TLV for chlorobenzene is 75 ppm. Benzene and even aliphatic hydrocarbons [A85] have been used as solvents in SPE cells. III.A.3.n Ethers Ethers, such as THF and 1,2-dimethoxyethane, have low dielectric constants (7.4 and 7.2, respectively), and the choice of supporting electrolyte is very limited. The ethers are difficult to reduce and are inert toward many metalorganic reagents [A36] that are soluble in the ethers. Ethers are thus suitable as a medium for the anodic addition of Grignard reagents to olefins [A223]. Dimethoxyethane and polyethylene glycol dimethyl ethers have the ability to solvate electrons [A224]. Dissociation constants for some acids and the relative strength of some bases in THF have been determined [A225].

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As a supporting electrolyte, NaClO4 or LiClO4 may be used in both solvents; TBAI in THF and TBAP in dimethoxyethane have also been employed. As a reference electrode, the Ag/Ag+ electrode functions satisfactorily [A36]. Ferrocene and bis(biphenyl)chromium have also been used as reference electrodes in THF [A226]. The ethers form peroxides. These can be removed by distillation from LiAlH4. They must be stored under nitrogen or distilled from LiAlH4 or Na directly into the electrolytic cell. If the water content of DME or THF is not negligible, the cathode may be passivated by a layer of a hydroxide formed during electrolysis, when an alkali metal salt is used as the supporting electrolyte [A227]. Purification of THF using sodium dianions of benzophenone and anthracene has been described [A228]. Protons and Lewis acids (EGA) can induce polymerization of THF; this may cause problems at the anode. The TLV for THF is 200 ppm. III.A.3.o Methylene Chloride and 1,1,2,2-Tetrachloroethane Electrolysis may be carried out in methylene chloride. It dielectric constant is low (ε = 9), and only the larger tetraalkylammonium salts are soluble in methylene chloride. Methylene chloride has been used for both anodic and cathodic reactions. The major advantages of using methylene chloride seem to be that certain cation radicals are more stable in methylene chloride than in the solvents usually employed in electrochemistry [A229,A230]. A disadvantage is that chloride ions produced at the cathode may diffuse to the anode compartment and interfere with the anodic reactions. This can partially be avoided by the addition of small amounts of acetic acid to the catholyte whereby the cathodic reaction becomes an evolution of hydrogen rather than formation of chloride ions. In electrosynthesis in systems of two immiscible liquids and a phase transfer catalyst, methylene chloride is in many cases preferred as the organic solvent. The addition of TFA and its anhydride (TFAn) to CH2Cl2 enhances the stability of cation radicals in the solution. A mixture of CH2Cl2–TFA–TFAn of 45:1:5 has been recommended [A49]. TLV is 100 ppm. 1,1,2,2-Tetrachloroethane has been found to be a good solvent for anodic investigations of fullerenes (Chapter 21). It may be purified by refluxing over CaH 2 and dried over P2O5 under high vacuum [A231]. III.A.3.p Sulfur Dioxide Sulfur dioxide is a solvent of low nucleophilicity and may be used when nucleophilic trapping by solvent is to be avoided. Its liquid range is somewhat inconvenient (−75 to −10°C), but its dielectric constant is reasonably high (17.6 at −20°C). The SO2 may be purified by washing with concentrated sulfuric acid, percolated through P2O5, and finally treated through a column of basic alumina [A232,A233]. As a supporting electrolyte, Bu4NClO4 or Pr4NPF6 may be used [A234–A236]; at 0.2 M solution of these salts, it has a conductivity of about 9 × 10 −3 Ω−l cm−1 at −22°C. CsAsF6 gives a potential window to about 5 V vs. SCE, enough to oxidize methane; however, the solubility (and thus conductivity) of CsAsF6 in SO2 is too low to allow a preparative oxidation [A232]; however, Bu4NBF4 has been used for preparative methylthiation [A237]. As a reference electrode, Ag/0.1 M AgNO3 in acetonitrile has been used. The anodic range of perchlorates or hexafluorophosphates in sulfur dioxide is large. AlCl3 may also be used as a supporting electrolyte [A238]; cation radicals are considerably stabilized in this medium.

III.A.4

SALTS

In certain electrolytic media, the concentration of salts is so high that the medium could be regarded as consisting of a solvated salt. This is thus the case in the strong solutions (10 M) of many salts in ammonia in which the vapor pressure of ammonia at ambient temperature is less

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than 1 atm. It also applies to the strong solutions of tetraalkylammonium p-toluenesulfonates in water, used sometimes in organic electrochemistry. A genuine liquid salt is tetrahexylammonium benzoate (THAB) [A40], which is a viscous liquid capable of acting as both solvent and supporting electrolyte; it has, however, a rather high resistance. Although it is a fairly polar solvent, THAB is miscible with benzene, toluene, and carbon tetrachloride; a mixture of THAB–toluene (75:25) has a rather low specific resistance of 3.8 kΩ and a much lower viscosity than pure THAB, which is about as viscous as glycerol. Relatively little water is miscible with THAB; addition of 1% by volume of water to THAB results in a two-phase system. THAB is soluble in acetone. THAB is useful for reductions; the cathodic limit at an Hg electrode is about −2.6 V versus Ag/AgCl; at a platinum electrode, the anodic limit is at about 0.3 V. THAB has been prepared by neutralizing tetrahexylammonium hydroxide, obtained from the iodide by treatment with Ag2O, with benzoic acid and evaporation of the water. The analysis of the dried product corresponds to a hemihydrate. The material is somewhat hygroscopic but can be dried by brief heating to 90°C. Voltammetry has also been made in (toluene)3tetrabutylammonium tetrafluoroborate; the resistance of the solution is between that of DMF and THF [A239] but closer to that of DMF. On addition of proton donors, a low “buffer capacity” of the solvent is observed. Aluminum chloride in mixture with an alkylpyridinium chloride, preferentially butylpyridinium chloride, has certain possibilities as a medium for studying electrochemically generated cation radicals. The 0.8:1 AlCl3–BuPy+Cl− mixture is molten at room temperature but has generally been used at 40°C [A240]. A mixture of l-methyl-3-ethylimidazolium chloride with aluminum chloride also gives an ionic liquid that may be used as electrolytic medium [A241]. The medium is not quite inert; for example, quinone reacts with an excess of chloride to 2-chlorohydroquinone in an acidcatalyzed Michael addition [A242]. As a reference electrode, an aluminum wire was used. l-Methyl-3-ethylimidazolium chloride is in MeCN reduced at −2.35 versus Ag in 0.1 M TBAP [A243]. At higher temperature, certain quaternary ammonium salts (40–150°C) [A244] may be used as media for electrolysis.

III.A.5

SUPERCRITICAL FLUIDS

SCFs have been used mainly for selective extraction of compounds; the solubility of a compound in a given solvent is in many cases vastly different under ambient and supercritical conditions. Thus, supercritical water dissolves both polar and nonpolar compounds, which may be explored in electrochemistry. When temperature and pressure approach the critical values, the internal structure of the solvent is loosened and the viscosity, the dielectric constant, and the density diminish; the dielectricity constant ε of water thus diminishes from 80 at 25°C to 5.2 at 647°C at 221 bar [A245]. Cyclic voltammetry in supercritical water-0.2 M NaHSO4 [A246] and ammonia-0.14 M CF3SO3K [A134,A246] of some organic compounds shows that this electroanalytical technique was applicable under these conditions. The behavior of phenazine in NH3 at −40°C and under supercritical conditions, for example, was analogous; two reversible reductions were found in both cases [A246]. Dimethyl carbonate has been prepared from CO and MeOH on anodic oxidation in a supercritical mixture of CO2 and MeOH [A247]. Supercritical carbon dioxide (scCO2) has a dielectric constant ε = 1.6 (cyclohexane ε = 2.0), and even though tetrakis(decyl)ammonium tetraphenylborate has some solubility in scCO2, the conductivity is too low for practical electrochemistry [A248]. 1,1,1,2-Tetrafluoroethane (HFC 134a) has a higher polarity in the supercritical state, and well-behaved CV may be obtained; with TBAClO4, a potential “window” of about 9 V has been obtained [A249].

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A202. JF Coetzee. Pure Appl Chem 49:211, 217, 1977; JF Coetzee. In: Recommended Methods for Purification of Solvents and Tests for Impurities (JF Coetzee, ed.) Oxford, U.K.: Pergamon, 1982, p. 16. A203. B Bry, B Tremillon. J Electroanal Chem 30:457, 1971; 46:71, 1973. A204. M Machtinger, MJ Vuaille, B Tremillon. J Electroanal Chem 83:273, 1977. A205. RE Johnson. NASA Tech Memo, NASA TMX-1283, 1966; NASA TMX-1380, 1966; Chem Abstr 66:108836, 1967. A206. K Chiba, M Fukuda, S Kim, Y Kitano, M Tada. J Org Chem 64:7654, 1999 and references therein. A207. AKR Unni, L Ellias, HI Schiff. J Phys Chem 67:1216, 1963. A208. LS Marcoux, JM Fritsch, RN Adams. J Am Chem Soc 89:5768, 1967. A209. J Courtot-Coupez, M L’Her. Bull Soc Chim Fr 1631, 1970. A210. DR Cogley, JC Synnott, JN Butler, E Grunwald. Electrochemical Society Meeting. Washington, DC, May 1971, Extended Abstract 122. A211. AN Dey, BP Sullivan. J Electrochem Soc 117:222, 1970. A212. G Eichinger. J Electroanal Chem 74:183, 1976. A213. G Eggert, J Heitbaum. Electrochim Acta 31:1443, 1986. A214. B Rasch, E Cattaneo, P Novak, W Vielstich. Electrochim Acta 36:1397, 1991; S Wasmus, W Vielstich. Electrochim Acta 38:541, 1993. A215. J Hlavaty, P Novak. Electrochim Acta 37:2595, 1992. A216. E Cattaneo, B Rasch, W Vielstich. J Appl Electrochem 21:885, 1991. A217. A Caiola, G Guy, JC Sohm. Electrochim Acta 15:555, 1970. A218. DP Boden, LM Mukherjee. Electrochim Acta 18:781, 1973. A219. FGK Baucke, CW Tobias. J Electrochem Soc 116:34, 1969. A220. T Fujinaga, K Izutsu. In: Recommended Methods for Purification of Solvents and Tests for Impurities (JF Coetzee, ed.) Oxford, U.K.: Pergamon, 1982, p. 19. A221. R Lines, VD Parker. Acta Chem Scand B31:369, 1977. A222. S Nakabayashi, A Fujishima, K Honda. J Electroanal Chem 111:391, 1980. A223. HJ Schäfer, Chem Ing Tech 42:164, 1970; HJ Schäfer, H Küntzel. Tetrahedron Lett 3333, 1970. A224. A Misono, T Osa, T Yakamichi, T Kodoma. J Electrochem Soc 115:266, 1968. A225. BK Deshmuhk, S Siddiqui, JF Coetzee. J Electrochem Soc 138:124, 1991. A226. SI Bailey, WP Leung, IM Ritchie. Electrochim Acta 30:861, 1985. A227. A Caillet, G Demange-Guerin. J Electroanal Chem 40:69, 1972. A228. F Paolucci, M Carano, P Ceroni, L Mottier, S Roffia. J Electrochem Soc 146:3357, 1999. A229. J Phelps, KSV Santhanam, AJ Bard. J Am Chem Soc 89:1752, 1967. A230. K Nyberg. Acta Chem Scand 24:1609, 1970. A231. Y Yang, F Arias, L Echegoyen, LPF Chibante, S Flanagan, A Robertson, LJ Wilson. J Am Chem Soc 117:7801, 1995. A232. E Garcia, J Kwak, AJ Bard. Inorg Chem 27:4377, 1988; E Garcia, AJ Bard. J Electrochem Soc 137:2752, 1990; JB Chlistunoff, AJ Bard. Inorg Chem 31:4582, 1992. A233. M Dietrich, J Heinze. J Am Chem Soc 112:5142, 1990. A234. LL Miller, EA Mayeda. J Am Chem Soc 92:5818, 1970. A235. P Castellonese, PC Lacaze. CR Acad Sci Ser C 274:2050, 1972. A236. LA Tinker, AJ Bard. J Am Chem Soc 101:2316, 1979. A237. RS Glass, VV Jouikov. Tetrahedron Lett 40:6357, 1999. A238. PC Lacaze, JE Dubois, M Delmar. J Electroanal Chem 102:135, 1979. A239. A Fry, J Touster. J Org Chem 51:3905, 1986. A240. J Robinson, RA Osteryoung. J Am Chem Soc 101:323, 1979; J Robinson, RA Osteryong. J Am Chem Soc 102:4415, 1980; RJ Gale, RA Osteryoung. J Electrochem Soc 127:2167, 1980. A241. M Lipsztajn, RA Osteryoung. J Electrochem Soc 130:1968, 1983; Electrochim Acta 29:1349, 1984; L Janiszewska, RA Osteryoung. J Electrochem Soc 134:2787, 1987; 135:116, 1988; TA Zawodzinski Jr, RT Carlin, RA Osteryoung. Anal Chem 59:2639, 1987; JF Oudard, RD Allendoerfer, RA Osteryoung. J Electroanal Chem 241:231, 1988. A242. FA Uribe, RA Osteryoung. J Electrochem Soc 135:378, 1988. A243. J Xie, TL Riechel. J Electrochem Soc 145:2660, 1998. A244. JE Gordon. J Am Chem Soc 87:4347, 1965. A245. M Uematsu, EU Franck. J Phys Chem Ref Data 9:1291, 1980. A246. WM Flarsheim, YM Tsou, I Trachtenberg, KP Johnston, AJ Bard. J Phys Chem 90:3857, 1986; RM Crooks, AJ Bard. J Phys Chem 91:1274, 1987; J Electroanal Chem 240:253, 1988; 243:117, 1988.

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A247. RA Dombro Jr, GA Prentice, MA McHugh. J Electrochem Soc 135:2219, 1988. A248. AP Abbott, JC Harper. J Chem Soc Faraday Trans 92:3895, 1996; AP Abbott, TA Claxton, J Fawcett, JC Harper. J Chem Soc Faraday Trans 92:1747, 1996; Proc Electrochem Soc 97–28:83, 1998. A249. AP Abbott, CA Eardley, JC Harper, EG Hope. J Electroanal Chem 457:1, 1998; AP Abbott, CA Eardley. J Phys Chem B 102:8574, 1998.

III.b

APPENDIX b*: ELECTROLyTES

The passage of electric current through a nonmetallic liquid is dependent on ions. The choice of electrolyte depends on such properties as its solubility, dissociation constant, mobility, discharge potential, and protic activity. It is desirable to employ a salt with high solubility, complete dissociation, high mobility of the ions, and a numerically high discharge potential. The role of adsorption of ions at the electrode is discussed elsewhere. The electrolyte may influence the rate constant for electron transfer from an electrode to a substrate.

III.B.1

ANIONS

The choice of anion is of most importance in anodic reactions in which the choice may influence the product distribution. Unless the oxidized anion is desired in an indirect electrolytic oxidation (Chapter 29)†, an anion that is oxidizable with difficulty, such as perchlorate, tetrafluoroborate, hexafluorophosphate, or nitrate, is chosen. In aqueous solution, the choice of anion is less critical than in nonaqueous solvents. In cathodic reactions, the choice of cation has highest priority, and the anion is generally selected on the basis of solubility of the salt or the ability of the anion to assist in dissolving the substrate. III.b.1.a Perchlorate The perchlorate ion has a high anodic discharge potential, and only in a few solvents, such as acetonitrile or nitromethane, is the perchlorate ion oxidized in preference to the solvent. When oxidized, the primarily formed perchlorate radical decomposes into oxygen and chlorodioxide radical [B1]. Caution must be exercised when handling perchlorates. Silver perchlorate is explosive; grinding it in a mortar should be avoided. For use in reference electrodes, it may be prepared by anodic dissolution of silver in a perchlorate medium [B2]. Evaporation of solutions of organic perchlorates is discouraged but, if necessary, must occur in vacuo and at moderate temperatures. Even sodium perchlorate, which can be dried at 100°C without undue risk, may cause explosions when heated with an organic solvent. Sodium perchlorate is soluble in many organic solvents. It is hygroscopic but can be dried at 110°C. Drying NaClO4, H2O often results in a sintered mass that must be broken and dried. By recrystallization of NaClO4 at temperatures about 80°C, the anhydrous salt can be obtained, which can be dried without difficulty. Lithium perchlorate is very hygroscopic. Tetraalkylammonium perchlorates are less hygroscopic, and some of them may be recrystallized from water; they are useful for voltammetric studies in organic solvents. III.b.1.b Tetrafluoroborate The discharge potential of the tetrafluoroborate ion is slightly more positive [B3] than that of the perchlorate ion. Tetrafluoroborates do not have a tendency to explode, and they are recommended instead of perchlorates for use in preparative work when the solvent must be removed during the workup. It has been reported that BF4− as anion in a supporting electrolyte in DMF gives in CV an anodic peak around 0.8 V versus SCE at a glassy carbon electrode [B4]. * The following text is taken from Lund H., in Lund H. and Hammerich O., Organic Electrochemistry, 4th edn., Marcel Dekker, New York, 2001, pp. 272–277; references have been renumbered, cross-references have been adapted, and refer to chapters of the present, fifth edition if not otherwise noted; some obvious (typographical) errors in the original have been corrected. † This chapter reference refers to the forth edition.

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Tetrabutylammonium tetrafluoroborate is only slightly soluble in water but is soluble in organic solvents. It is prepared by mixing solutions of NaBF4 with a soluble tetrabutylammonium salt. Its solubility properties are convenient for its use in preparative electrolysis. During the workup, the organic solvent may be diluted with water, whereby the salt precipitates are recovered for future use. It may be recrystallized from toluene; the solution is dried by azeotropic distillation (or by molecular sieves A4); the compound crystallizes slowly on cooling [B5]. III.b.1.c Hexafluorophosphate, Hexafluoroarsenate Hexafluorophosphates are slightly more resistant toward anodic oxidation than tetrafluoroborates and perchlorates [B3]; in a 0.1 M solution of these salts in acetonitrile, the potentials (versus Ag/0.l M Ag+ electrode) at a platinum disk electrode (0.12 cm2) at 1 mA were for perchlorate, 2.48 V; for tetrafluoroborate, 2.91 V; and for hexafluorophosphate, 3.02 V [B3]. Hexafluoroarsenates have been used in liquid SO2 [B6,B7]. III.b.1.d Trifluoromethanesulfonate The trifluoromethanesulfonates have certain advantages over perchlorates (nonexplosive) and tetrafluoroborates (even less reactive toward oxidizing agents) [B8]. The tetraalkylammonium trifluoromethanesulfonates are just as soluble or more soluble in the commonly used organic solvents than the corresponding perchlorates or fluoroborates. They can be prepared either by metathesis or by alkylation of a tertiary amine by an ester of trifluoromethanesulfonic acid. III.b.1.e Nitrate Nitrate ion is not as resistant to oxidation as perchlorate or tetrafluoroborate; in nitromethane [B9] solution, it is oxidized anodically to NO2+ and oxygen, and the NO2+ reacts either with the water always present in “dry” organic solvents to form nitric acid or with nitrate ion to form N2O5. Nitrates, to a higher degree than perchlorates [B10], form ion pairs in some solvents. Tetrabutylammonium nitrate can be prepared conveniently by mixing an aqueous solution of NaNO3 with the commercially available tetrabutylammonium hydrogen sulfate. The salt is then extracted into methylene chloride, giving an almost quantitative yield of Bu4NNO3. III.b.1.f Aromatic Sulfonates The use of aromatic sulfonates stems from their ability to assist the solution of organic substrates in an aqueous medium. This “hydrotropic” effect [B11–B13] is accentuated if a quaternary ammonium sulfonate, rather than an alkali metal sulfonate, is used. The aromatic sulfonates have been employed mostly for reductions, but they would be expected to be reasonably stable toward oxidation. To obtain a reasonable hydrotropic effect, rather high concentrations are used; in the original hydrodimerization reaction of acrylonitrile, the medium may be regarded as either a solution of tetraalkylammonium p-toluenesulfonate in water or the salt containing some water. III.b.1.g Carboxylate Ion In the Kolbe reaction and in anodic carboxylation reactions, salts of carboxylic acids function as both substrate and electrolyte. Usually the presence of other anions diminishes the yield of the Kolbe reaction, whereas the presence of anions that are oxidizable with difficulty, like bicarbonate or perchlorate ions, favors the related Hofer–Moest reaction. For reductions in aprotic media, tetraalkylammonium salts of derivatives of oxalic and formic acid may be employed; in many cases, the reaction may then be performed in an undivided cell, as the reaction at the counter electrode (oxidation of the anion to carbon dioxide) generally does not interfere with the cathodic reaction [B14].

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III.b.1.h Tetramethylaluminate and Tetraphenylborate The sodium salts of these anions are soluble in THF and have been found suitable as supporting electrolyte in this solvent [B15,B16].

III.B.2 CATIONS The discharge potential of a cation determines whether it is useful; so in practice, only alkali and alkali earth metal ions together with ammonium and tetraalkylammonium ions are used. The nondischarge of a cation is a sine qua non, but in addition, a choice of cation must take into consideration ion pair formation, solvation, and adsorption of particular species (see, e.g., References B17–B19). III.b.2.a Lithium Ions Many lithium salts, such as lithium perchlorate and the halides, are soluble in nonaqueous solvents. The reduction potential of Li+ depends on the electrode and the solvent. At a mercury cathode, amalgam formation takes place, whereas formation of lithium metal occurs at platinum in aprotic media. Lithium metal is less reactive than sodium, and in some solvents, sodium attacks the solvent, whereas lithium is unreactive. A small water content in an aprotic solvent may react with lithium (or Li+ may react with hydroxyl ions formed at the cathode) to form lithium hydroxide, which may cover the electrode with an insoluble, insulating layer. III.b.2.b Sodium Ions In aqueous buffers, sodium ion is a common cation, but in nonaqueous solvents, practically only NaClO4 is useful. Even NaBF4 is not suitable in most aprotic solvents. III.b.2.c Magnesium Ions Magnesium ions interact with radical ions in aprotic media, which may influence the course of the reaction; in some cases, Mg+ also stabilizes the product and thereby helps avoiding unwanted further reactions [B20,B21]. III.b.2.d Tetraalkylammonium Ions The most commonly used quaternary ammonium salts are tetrabutylammoniumperchlorate (TBAP), tetrafluoroborate (TBAT), the halides (TBACl, TBAB, and TBAI), and the corresponding tetraethylammonium salts, such as the perchlorate (TEAP), but also the tetramethyl- or tetrapropylammonium salts have been employed; the former cannot undergo a base-promoted Hofmann elimination. However, evidence has been found for the formation of trimethylammonium methylide [B22]. In nonpolar solvents, it may be necessary to employ tetrahexyl- or tetraoctylammonium salts. The tetraalkylammonium ions are soluble in many nonaqueous media, and they may be extracted from an aqueous solution by means of chloroform or methylene chloride [B23,B24], and tetraalkylammonium salts may thus be prepared by ion extraction [B24]. Tetrakis(decyl)ammonium tetraphenylborate is soluble even in hexane [B25,B26]. Reduction of R4N+ at a glassy carbon electrode in DMF yields an alkyl carbanion, which deprotonates residual water or R4N+ in a Hofmann elimination with formation of an alkene and a trialkylamine. Tetraethylammonium is attacked more easily than tetrabutylammonium [B27]. The discharge potential becomes slightly more negative with increasing size of the ion, but in practice, there is not much gain by going beyond tetrabutylammonium ions [B28]. The reduction route of these ions is discussed elsewhere (Chapter 16). In aqueous solution, the tetraalkylammonium ions are adsorbed at the electrode, where they form a layer with a low proton activity. This property is used in the commercial production of

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adiponitrile from acrylonitrile (Chapter 31)*. The adsorption of the ions depends on the potential of the electrode, the electrode material, the composition of the medium, and the size of the cation. The rate constant for electron transfer from an electrode to a substrate decreases with the size of the cation; for instance, the rate constants for electron transfer at a mercury electrode in MeCN to a substrate were always smaller when tetra-n-heptylammonium perchlorate was the electrolyte than when tetraethylammonium perchlorate was used [B29]. Quaternary ammonium salts are readily accessible; some are commercially available, and their preparation presents no problems [B23,B24]. Reducible impurities may be removed by electrolysis at a sufficiently negative potential. The production of tetraalkylammonium hydroxide using an anion [B30] or cation [B31] exchange membrane has been reported. III.b.2.e Sulfonium Salts Tris(dimethylamino)sulfonium tetrafluoroborate (TASBF4) has been investigated as a supporting electrolyte in MeCN and CH2C12. Compared with TBABF4, the conductance of TASBF4 in CH2C12 is slightly higher, suggesting a higher degree of dissociation of the latter [B32]. III.b.2.f Cryptates Crown ethers and other cryptates form stable complexes with several metal cations, which generally are soluble in nonaqueous media, even in rather nonpolar solvents [B33]. In some cases, such complexes may be considered supporting electrolytes, especially when some of the special properties of a given cryptate are desirable also in other contexts. By complexing Ag+ with a cryptate, the Ag/AgCryp+ electrode becomes stable in DMF [B34]. III.b.2.g Polyelectrolytes Polyelectrolytes having ion exchange properties may be coated on one or both electrodes [B35–B40], which makes it possible to use virtually nonconducting liquids. This was discussed in Section II.D.3. A copolymer of 4-vinylpyridine or styrene with a diaryl-4-(l-trifluoromethylvinyl)-aryl amine has sufficiently solubility and conductivity to serve as both supporting electrolyte and redox mediator; after use, it can be recovered by ultrafiltration [B41].

III.B.3

BUFFERS

Control of an electrolytic reaction often requires that the proton activity remains within acceptable limits during the electrolysis. For small-scale electrolytic preparations (less than about 10 g/liter of substrate), a sufficiently high initial concentration of buffer, acid, or base is adequate in aqueous solution; for large-scale electrolysis, a controlled addition of protons during a reduction must be provided. This addition may be controlled by a pH-stat or coupled to the current integrator. In aprotic media, a proton donor, electrophile, or nucleophile may play a similar role as buffers in aqueous media. III.b.3.a Acids Hydrochloric acid is a convenient acid for laboratory-scale electrolysis; it is easy to remove during the workup by evaporation. The oxygen-containing acids, such as sulfuric or perchloric acid, must be neutralized before basic products can be extracted. Sulfuric acid is most conveniently neutralized by concentrated ammonia rather than by sodium hydroxide, whereby the precipitation of sodium sulfate is avoided. Even in acetic acid, perchloric acid is a strong acid. Anhydrous solutions are made by adding acetic anhydride equivalent to the water present in the acetic acid and introduced together with the perchloric acid. An excess of acetic anhydride should be avoided as this solution is a very strong * This chapter reference refers to the fourth edition.

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acylating agent, and it forms yellow condensation products on standing. Anhydrous perchloric acid in MeCN has been prepared by oxidation of H2 to H+ at a platinized platinum electrode in a solution of TBA+ClO4− in MeCN [B42]. In aprotic media, the addition of proton donors may be important. If the proton donor is too strong, protons may be reduced in preference to the substrate and a too weak proton donor may not furnish sufficient protons. Often a phenol is an acceptable compromise. Extreme acidic conditions may be obtained in superacids, such as SbF5 or BF3 in HF or SbF5 in FSO3H. III.b.3.b bases Reductions in aqueous alkaline solution do not present many problems; the choice of cation was discussed earlier. In nonaqueous solvents, bases may be generated electrolytically (Chapter 43) [B43,B44] or formed by reaction of the solvent with an alkali metal. The conjugate base of the solvent is the strongest base obtainable in any medium. During oxidation in a nonaqueous solvent, the removal of a proton or an attack of a nucleophile on the oxidized substrate may be a chemical step in an ECE reaction. To explore which reaction is operating, one can, in parallel voltammetric experiments, use two compounds that are equal in base strength but differ widely in nucleophilicity; such a pair is pyridine and 2,6-dimethylpyridine. III.b.3.c buffer Systems The use of buffers in electrolysis to maintain a desired pH warrants some consideration. During a reduction, protons are consumed at the cathode surface, and unless they are replenished rapidly, the pH in the reaction layer is considerably higher than in the bulk of the solution. A high concentration of a buffer system, which exchanges protons rapidly, is required to maintain a desired pH. In aqueous solution, phosphate, borate, and carboxylate buffers are applicable, but in mixed solvents, the solubility of phosphate and borate buffers is low; in such mixtures, carboxylate buffers may be used in slightly acid medium and amine buffers in slightly alkaline medium. To avoid excessive buffer concentrations, it is advisable to maintain pH in the bulk of the solution constant by means of a pH-stat. Buffering in an aprotic medium is sometimes a problem. Proton donors, such as a suitable phenol, malonic ester, or amine salt, together with the corresponding base, may be an acceptable solution. Guanidine perchlorate [B45] has been proposed as an efficient proton donor (and supporting electrolyte) in aprotic solvents, as it brings the protons into the Helmholtz double layer.

REFERENCES FOR APPENDIX b B1. B2. B3. B4. B5. B6. B7. B8. B9. B10. B11. B12. B13. B14. B15. B16.

BS Jensen, VD Parker. J Am Chem Soc 97:5211, 1975. G Cauquis, D Serve. Bull Soc Chem Fr 302, 1966. M Fleischmann, D Pletcher. Tetrahedron Lett 6255, 1968. M Chandrasekaran, M Noel, V Krishnan. Talanta 37:695, 1990. H Lund. Unpublished observation, 1986. LA Tinker, AJ Bard. J Am Chem Soc 101:2316, 1979. RS Glass, VV Jouikov. Tetrahedron Lett 40:6357, 1999. K Rousseau, GC Farrington, D Dolphin. J Org Chem 37:3968, 1972. G Cauquis, D Serve. CR Acad Sci Ser C 262:1516, 1966. H Yeager, B Kratochvil. J Phys Chem 73:1963, 1969. Y Kargin, LM Vorontsova, LZ Manapova, SV Kuzovenko. Elektrokhimiya 23:415, 1987. RH McKee. Ind Eng Chem 38:382, 1946. MM Baizer. J Electrochem Soc 111:215, 1964. R Engels, WJM van Tilborg, CJ Smit. Sandbjerg Meeting. 1981, Abstracts of Papers, p. 73. BL Funt, D Richardson, SN Bhadani. Can J Chem 44:711, 1966. N Yamazaki, S Nakaham, S Kambara. Polym Lett 3:57, 1965.

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330 B17. B18. B19. B20. B21. B22. B23. B24. B25. B26. B27. B28. B29. B30. B31. B32. B33. B34. B35. B36. B37. B38. B39. B40. B41. B42. B43. B44. B45.

Organic Electrochemistry T Fujinaga, K Izutsu, T Nomura. J Electroanal Chem 29:203, 1971. MM Baizer, JP Petrovich. J Electrochem Soc 114:1023, 1967. JP Petrovich, MM Baizer. J Electrochem Soc 118:447, 1971. D Pletcher, L Slevin. J Chem Soc Perkin 2217, 1996; 2005, 1995. S Kashimura, Y Murai, M Ishifune, H Masuda, H Murase, T Shono. Tetrahedron Lett 36:4805, 1995. KL Vieira, MS Mubarak, DG Peters. J Am Chem Soc 106:5372, 1984. A Brändström, K Gustavii. Acta Chem Scand 23:1215, 1969; A Brändström, P Berntsson, S Carlsson, A Djurhuus, K Gustavii, U Junggren, B Lamm, B Samuelsson. Acta Chem Scand 23:2202, 1969; A Brändström. Preparative Ion Pair Extraction. Stockholm, Sweden: Apotekersocieteten, 1974. M Makosza, E Bialecka. Synth Commun 6:313, 1976. AP Abbott, JC Harper. J Chem Soc Faraday Trans 92:3895, 1996; AP Abbott, TA Claxton, J Fawcett, JC Harper. J Chem Soc Faraday Trans 92:1747, 1996; Proc Electrochem Soc 97–28:83, 1998. AP Abbott, CA Eardley, JC Harper, EG Hope. J Electroanal Chem 457:1, 1998; AP Abbott, CA Eardley. J Phys Chem B 102:8574, 1998. CE Dahm, DG Peters. J Electroanal Chem 402:91, 1996. HO House, E Feng, NP Peet. J Org Chem 36:2371, 1971. RA Peterson, DH Evans. J Electroanal Chem 222:129, 1987, and references therein. JR Ochoa Gomez, M Tarancon Estrada. J Appl Electrochem 21:354, 1991. M Kashiwase, H Harada, K Tomiie. In: Recent Advances in Electroorganic Synthesis (S Torii, ed.) New York: Elsevier, 1987, p. 467. JY Becker, T Berzins, BE Smart, T Fukunaga. J Electroanal Chem 248:363, 1988. JF Coetzee, JM Simon, RJ Bertozzi. Anal Chem 41:766, 1969. K Igutsu, M Ito, E Sarai. Anal Sci 1:341, 1985. F Beck, H Guthke. Chem Ing Tech 41:943, 1969. Z Ogumi, K Nishio, S Yoshizawa. Electrochim Acta 26:1779, 1981; Z Ogumi, H Yamashita, K Nishio. Electrochim Acta 28:1687, 1983. J Sarrazin, A Tallec. J Electroanal Chem 137:183, 1982; E Raoult, J Sarrazin, A Tallec. J Appl Electrochem 14:639, 1984; 15:85, 1985; Bull Soc Chem Fr 1200, 1985; J Membr Soc 30:23, 1987. Z Ogumi, M Inaba, SI Ohashi, M Uchida, ZI Takehara. Electrochim Acta 33:365, 1988. M Inaba, Z Ogumi, ZI Takehara. J Electrochem Soc 140:19, 1993. VA Grinberg, VN Zhuravleva, YB Vasil’ev, VE Kazarinov. Elektrokhimiya 19:1447, 1983. R Wend, E Steckhan. Electrochim Acta 42:2027, 1997. K Pekmez, M Can, A Yildiz. Electrochim Acta 38:607, 1993. PE Iversen, H Lund. Tetrahedron Lett 3523, 1969. MM Baizer, JL Chruma, DA White. Tetrahedron Lett 5209, 1973; RC Hallcher, MM Baizer. Liebigs Ann Chem 737, 1977. R Breslow, RF Drury. J Am Chem Soc 96:4702, 1974.

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8

Application of Ionic Liquids, Emulsions, Sonication, and Microwave Assistance John D. Watkins and Frank Marken

CONTENTS I. Introduction: The Need for Special Conditions in Organic Electrosynthesis ...................... 331 II. Ionic Liquids in Organic Electrosynthesis ........................................................................... 332 III. Supercritical Conditions in Organic Electrosynthesis .......................................................... 335 IV. Emulsions and Biphasic Conditions in Organic Electrosynthesis ........................................ 336 V. Ultrasound in Organic Electrosynthesis ............................................................................... 339 VI. Microwaves in Organic Electrosynthesis .............................................................................340 VII. Special Conditions in Specific Applications ........................................................................ 341 References ...................................................................................................................................... 341

I.

INTRODUCTION: THE NEED FOR SPECIAL CONDITIONS IN ORgANIC ELECTROSyNTHESIS

Electrochemical processes are predominantly heterogeneous in nature requiring mass transport and in some cases thermal reaction control. The need for special conditions in electrosynthesis is therefore intertwined with the control of the mechanism and linked either to the enhancement of reaction selectivity, conversion rate, or product yield. Increasingly however, it is also the search for green conditions that drives the exploration of special conditions [1]. For example, the use of excess solvents has long been regarded as detrimental to the environment, with an onus to reduce their quantity or replace them with less hazardous alternatives. It is for this reason that supercritical fluids and various types of room temperature ionic liquids (RTILs) are receiving much attention in standard organic preparations as well as electrosynthesis. Electrosynthesis often allows access to molecules inaccessible by standard synthetic procedures [2], but the need for special electrosynthesis equipment often limits its use to niche applications. Product separation from electrolytes can be time consuming and adds an extra separation step not usually required in organic synthesis, so either the electrochemical approach must be of sufficient benefit to outweigh this drawback or new methods to avoid this problem are required. For example, emulsion-based synthesis [3] or biphasic methods [4] have been proposed to eliminate the separation step by keeping electrolytes and reagents separated from products. The additional problem of beneficially using the often-ignored counterelectrode process can be approached, for example, by paired electrosynthesis conditions [5] or by the choice of a very small counterelectrode in a singlecompartment reactor to push the counterelectrode process potential into the solvent decomposition region with minimal loss of product [6]. Generally, new methods for improving the rate, selectivity, and yield in organic electrochemistry are desirable. It is for this reason that sonication [7] and microwave [8] techniques are sometimes considered. These techniques are readily focused on the working electrode surface to create 331

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synthesis conditions that are more effective than conventional stirring and heating techniques. In combination with the application of special reaction media, for example, ionic liquids or emulsions, entirely new processes may in future be possible. The following sections summarize some of the recent progress for methods that could be of benefit in electroorganic synthesis.

II. IONIC LIQUIDS IN ORgANIC ELECTROSyNTHESIS Ionic liquids are of special interest to the electrochemist due to their inherent ionic conductivity, tunable physical properties, low volatility, huge range of structural diversity, and potential for reuse [9–12]. RTILs (see Table 8.1) are most widely studied and are defined as an ion pair material that is sufficiently disordered at, or near to, room temperature so as to exist as a liquid [13–15]. To create TAbLE 8.1 Examples for RTIL Molecular Structures. (a) RTIL cations, (b) RTIL anions, and (c) Distillable Ionic Liquids

+ N R N

+ N

CH3

+ N

R

R

CH3

1-alkyl-3-methylimidazolium [CnC1im]+

1-alkylpyridimium [Cnpyr]+

1-alkyl-1-methylpyrrolidinium [CnC1pyrr]+

R4N+

R4P+

R3S+

Tetraalkylammonium [C4N]+

Tetraalkylphosphonium [C4P]+

Trialkylsulfonium [C3S]+

(a)

O F3C

S

– N

O

O–

O S

CF3

S

F3C

O

Bis(trifluoromethylsufonyl)imide [NTf2]–

NC – N

O

O Trifluoromethanesulfonate triflate, [OTf ]–

Alkylsulfate [CnSO4]–

F– F F

CN

Dicyanamide [N(CN)2]–

P

F– F F

F Hexafluorophosphate [PF6]–

(b)

CO2 + R1R2NH

R1R2NH – COOH

R1R2NH – COOH + R1R2NH

[R1R2NH2] +[R1R2NH – COO]–

(c) Source: Hallett, J.P. and Welton, T., Chem. Rev., 111, 3508, 2011.

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RSO4–

B F

F F

Tetrafluoroborate [BF4]–

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this state, mismatched ion pairs of low symmetry are often used. Generally, this is achieved by bulky asymmetric organic cations accompanied by almost any anion. Small halide-based anions generally lead to higher-temperature ionic liquids than inorganic complex anions such as PF6 −. Furthermore, the tunability of physical properties allows parameters such as organic molecule solubility and water miscibility to be controlled. The physical properties of RTILs as well as their reactivity for electrochemistry have been extensively reviewed [16,17], and a new classification of solvation properties has been developed [18]. Pioneering early work in 1980 by Osteryoung and coworkers focused on electrochemistry in alkyl-pyridinium aluminates [19], and special cases have been reported such as N,N-dialkyl-ammonium-N′,N′-dialkylcarbamate distillable ionic liquids [20] and purely natural ionic liquids (or deep Eutectic mixtures) based on choline chloride [21]. In this chapter, some key electrochemical processes will be discussed with an emphasis on electrosynthetic transformations as well as some useful publications concerned with the general use of ionic liquids in electrochemistry. Although ionic liquids are known to be inherently ion conducting, due to their ionic structure, it is often their higher viscosity that limits their use as an effective medium for electrosynthesis. This higher viscosity is linked to poor mass transport and limited ionic conductivities compared to conventional organic solvent media [22,23]. Like most physical properties of RTILs, the viscosity can be finely tuned by the choice of the cation and anion, with long-chain alkyl groups on the cation giving an increased viscosity. To counter this problem, cosolvents have been employed [24] as well as the inclusion of dissolved gasses, which have been shown to significantly affect the rate of diffusion [25]. Initial studies have often been focused on electropolymerization processes [26–29]. However, many electrosynthetic transformations have now been reported in RTIL media, including Michael addition [30], oxidation [31–33], reduction [34–36], aromatic substitution [37], cyclization [38], fluorination [39–41], dehalogenation [42], and coupling reactions [36,43–45]. The Michael addition, reported by Palombi [30] (see Figure 8.1), uses a divided flow cell and a paired electrosynthetic approach where both the anodic and cathodic compartments contribute to the process. The Michael product at the cathode is derived from enolate formation, and the process at the anode is catalyzed by BF3 Lewis acid formed in situ from the ionic liquid. With a paired convergent process strategy, it was possible to improve the current efficiency. O BF–4

Xc 1a 1 mmol O–

EMImBF4 O

Pt anode

Pt cathode

O

EMImBF4 BF3

Xc O

O

O

O

O

Xc=

Xc’ O

1 mmol

N Bn

Cathodic current yield 7.3 mol of 2a per mol of charge

O

O

O

Xc

Anodic current yield 6.8 mol of 2a per mol of charge

2a Overall isolated yield 71%

FIgURE 8.1 Schematic of the proposed mechanism for a divided cell paired electrolysis method for a Michael addition in 1-ethyl-3-methyl-imidazolium tetrafluoroborate. (From Palombi, L., Electrochim. Acta, 56, 7442, 2011.) The same, approximately 70% diastereoisomeric ratio, was observed for anodic and cathodic cells.

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The Shono oxidation in ionic liquid media with the inclusion of methanol was investigated by Bornemann and Handy [24], and good result and recovery of ionic liquids in the presence of a sufficiently high level of methanol were observed. Barhdadi et al. have demonstrated some examples of electrolysis in RTILs including the 2,2′,6,6′-tetramethylpiperidin-N-oxyl (TEMPO)mediated oxidation of alcohols [31] and the nickel-catalyzed electroreductive coupling of organohalides [45]. In the first case, a highly efficient TEMPO-mediated alcohol oxidation was achieved with the viscosity of RTIL overcome by the addition of appropriate amounts of alcohol and base. The current efficiencies achieved were up to 100% with conversions also 100%, although enolizable products were not tolerated to this degree. The second case demonstrates the first steps toward an efficient coupling reaction of organic halides with each other or with activated alkenes via a nickel catalyst (see Table 8.2). Experiments were conducted with a sacrificial Mg or Al anode and a bipyridyl-NiBr 2 catalyst in 1-octyl-3-methyl-imidazolium tetrafluoroborate. Similar nickel coupling reactions have been demonstrated by Mellah et al., who show that the homocoupling of aryl halides can be achieved and note that upon separation the nickel catalyst remains in the RTIL while organic materials are washed out. The same RTIL phase can be reused for further reactions. Mellah et al. [44] have also demonstrated an oxidative aromatic coupling process in RTILs. Another electroreductive coupling reaction by Duran Pachon et al. [46] has shown another benefit of an RTIL-based synthesis in the stabilization of in situ palladium nanoparticles. In this reaction, a counterelectrode reaction–generated Pd nanoparticle cluster was used as a ligand-less catalyst for the Ullmann coupling reaction of aryl iodides in good yield and functional group tolerance.

TAbLE 8.2 Aryl Coupling and Addition Processes Catalyzed by Ni(0) Complexes in 1-Octyl-3-Methyl-Imidazolium Tetrafluoroborate 2ArX + 2e–

RTIL, cat NiBr2bpy

Ar – Ar + 2X–

Anode Fe or stainless steel

R–X C8H17Br Ph–CH2Br Ph–CH2Cl Ph–CCl2–Ph

Isolated R–R

yield (%)

C16H34 Ph–CH2–CH2–Ph Ph–CH2–CH2–Ph Ph2C = CPh2

48 75 78 68 R

RTIL, cat NiBr2

Z

ArX + R–CH = CH–Z Stainless steel anode Arbr Ph—Br Ph—Br Ph—Br

Activated olefin CH2 = CH–CO–Me CH2 = CH–CO2–Bu MeO2C

Ar

Product and isolated Ph–CH2–CH2–CO–Me Ph–CH2–CH2–CO–Bu CO2Me

CO2Me

yield (%) 58 61 41

Ph CO2Me 3–MeO–C6H4Br

Ch2= CH–CO–Me

3–MeO–C6H4–(CH2)2–CO–Me

Source: Barhdadi, R. et al., Chem. Commun., 1434, 2003.

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335

O O + CO2 (1 atm) R

–2.4 V vs. Ag/AgCl

O

O

Ionic liquids, rt R

FIgURE 8.2 Cathodic activation of epoxides to react with CO2 in ionic liquid media with a sacrificial anode (Mg, Al) and a Cu cathode. Less than 1 mol equivalent of charge was passed.

Other uses of ionic liquids in electrosynthesis have been proposed for the in situ generation of reactive intermediates. It has been found by Martiz et al. that in a highly fluorinated RTIL, dioxygen is highly soluble and easily reduced to form a stable superoxide [28]. Hydrogen peroxide has also been successfully generated in situ by the electroreduction of oxygen in RTILs by both Tang et al. [47] and Ho et al. [48]. The former of these publications has shown that hydrogen peroxide was created in an RTIL water-mixed phase under basic conditions for direct epoxidation of electrophilic alkenes while being reusable for multiple syntheses. Ho et al. have used a similar oxidation system for a wider range of alkene epoxidations. This has been achieved by using carbon dioxide with the electrogenerated HO2− to create the peroxymonocarbonate ion (HCO4 −) intermediate. This intermediate, with a manganese (MnSO4) catalyst, was able to epoxidize a much wider range of alkenes. For the reuse of the RTIL phase, the manganese had to be extracted. The facile cathodic formation of peroxodicarbonate from O2 and CO2 in dry ionic liquids has also been reported [49]. A fixation of carbon dioxide by epoxides based on RTIL electrochemistry has been reported by Yang et al. [50] in which the reaction results in cyclic carbamates (see Figure 8.2). Dehalogenation reactions carried out with a vitamin B12 catalyst are common in organic electrosynthesis, but it has been found that these reactions are further enhanced by the polarity increase in RTILs [42,51]. In conclusion, there is a lot more scope for beneficial applications of RTILs in electroorganic reactions [52].

III.

SUPERCRITICAL CONDITIONS IN ORgANIC ELECTROSyNTHESIS

Solvents such as ionic liquids, supercritical carbon dioxide, or water are sometimes classed as neoteric solvents [53]. In contrast to RTIL media, supercritical fluids provide low polarity with fast diffusion and often require elevated temperature conditions. Applications of supercritical CO2, for example, in heterogeneous catalysis have been highly successful [54]. Potential for supercritical electrosynthesis [55] has been proposed in particular for CO2-based media [56,57] and for carboxylation processes [58]. A supercritical fluid is a phase that is subjected to pressures and or temperatures above their critical point values, thus creating a new gas–liquid fluidlike state. A commonly used supercritical liquid is carbon dioxide (scCO2), which may be considered both environmentally friendly and nonhazardous as well as useful under relatively mild conditions (TC = 31.1°C, PC = 73.8 bar [59]). A potential problem limiting the use of scCO2 in electrosynthetic applications is its low polarity and thus incompatibility with ionic electrolytes. This problem has been addressed by the use of appropriate cosolvents such as acetonitrile [60]. The cosolvent approach has been used for the electrocarboxylation of α-chloroethylbenzene with the addition of an ionic liquid (N,N-diethylN-methyl-N-(2-methoxyethyl)ammonium bis-(tri-fluoromethanesulfonyl)amide) and as a function of applied CO2 pressure [61] (see Figure 8.3). A similar approach to circumvent the problem of electrolyte solubility in scCO2 has been demonstrated for the potentially high-yielding oxidation of benzyl alcohol to benzaldehyde in a mixed scCO2 ionic liquid mixed system [62]. Ionic liquid was used as an electrolyte, and scCO2 pressure was used to control the product yield. Facile product isolation was possible in scCO2. Furthermore, the ionic liquid phase was shown to be reused over many reactions and thus was waste-free.

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Organic Electrochemistry Cathode reaction Cl

+2e–

CO2

COO–

–

–Cl Anode reaction

Mg Bulk reaction R–COO– + Cl– + Mg2+

Mg2+

+

2e–

R-COOMgCl and/or (R-COO)2Mg H3O+

COOH

FIgURE 8.3 Overall reaction scheme proposed for the reduction of α-chloroethylbenzene in ionic liquid/ CO2 with a platinum cathode and sacrificial Mg anode. (From Hiejima, Y. et al., Phys. Chem. Chem. Phys., 12, 1953, 2010.)

Supercritical fluids may be considered true designer solvents in a similar way as ionic liquids. The key difference is that instead of a laborious synthetic approach to change chemical compositions, temperature and pressure and cosolvent may be varied to continuously change the solvent properties. The miscibilities of cosolvents such as ionic liquids can be varied such that the reaction is completed in a monophasic system, and with a change in pressure, separation is carried out in a biphasic system. The presence of a biphasic system, and in particular the presence of a separate liquid phase deposit on the working electrode, has to be considered as a potential source of artifacts in electrochemical measurements in supercritical media. Films of poly(ethylene oxide) at the electrode surface have been employed intentionally to conduct electrosynthetic transformation in supercritical CO2 [63].

IV. EMULSIONS AND bIPHASIC CONDITIONS IN ORgANIC ELECTROSyNTHESIS Emulsion or biphasic media are of significant importance in electrosynthesis [64,65] and have been employed in large-scale processes such as the Monsanto production of Nylon 66 [66]. It is possible to form emulsions from a range of phase combinations, but aqueous phase systems are most commonly employed. Water is an ideal green electrosynthetic solvent due to its high dielectric constant with the only drawback being its incompatibility with some organic functional groups or reaction intermediates. Biphasic systems allow the conductivity of a fully supported electrolytic phase to be utilized alongside a secondary nonpolar organic phase in which a capture reagent can be used to contain products and to allow easy separation as well as preventing overelectrolysis. An emulsion system acts to separate electrolyte and products so as to form an in situ extraction procedure. Direct interaction of nonpolar reactants in aqueous-based biphasic processes can be enhanced, for example, by power ultrasound [67], microwave [68], or high-shear ULTRA-TURRAX homogenization [69]. However, more often, emulsion electrosynthesis requires an indirect electrosynthetic strategy using mediators, as reviewed by Ogibin et al. [70]. The topic of electrosynthesis in microemulsion has been reviewed also by Rusling et al. [3,71] in 1997 emphasizing the use of surfactants as stabilizing groups for microemulsions with biphasic mediators. Biphasic mediators are redox species that are able to transfer charge across a phase boundary by shuttling electrons from the

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B12r Oil — water interface

–2Br–

B12s

R

R΄

R

B12r

B12s

R΄

Br +e–

Br 2–10 μm (a) Polar

Polar H

H

H

O

O

O H

C12H25 H Br

Nonpolar

H

H

O H

H

Br

Br

Br

H

θ H

CH3

H

CH3

C12H25

Nonpolar

(b)

FIgURE 8.4 (a) Reaction scheme for vitamin B12–mediated reductive dehalogenation of vicinal dibromides at the oil–water interface for a microdroplet of oil immobilized at an electrode surface. (b) Close-up scheme suggesting a rotational polarization of dibromides at the oil–water interface. (From Davies, T.J. et al., ChemPhysChem, 6, 2633, 2005.)

electrode and into another phase. A common mediator that has been used previously in singlephase electrosynthesis is vitamin B12, which has been shown to be also an effective biphasic mediator. In the work by Davies et al. [72], the mechanism of the vitamin B12 biphasic mediation was investigated, and it was shown that the reduced B12 is able to diffuse to the phase interface of an aqueous–organic boundary and reduce vicinal dibromide species to form olefins (see Figure 8.4). A comparison of homogeneous and heterogeneous chemical rate constants for dehalogenation of different dibromides suggested that the liquid–liquid interface does play an active role in this process. Periodate has also been shown to be an effective biphasic mediator system in electrosynthesis. Khan et al. [73] applied an iodate aqueous phase to in situ create the powerful and more hydrophobic periodate oxidizing agent, which was able to transfer to the organic phase aided by a phase transfer catalyst to oxidize organic species. The emulsion-mediated approach has also been applied to the sharpless asymmetric dihydroxylation of alkenes [74]. In this case, a mixed cyclohexane– water biphasic system was employed with a ferricyanide mediator to maintain the concentration of an osmium(VIII) catalyst. The organic reaction is kept separate in the cyclohexane phase for ease of work-up. The use of a chiral oxaziridinium salt as organocatalyst for epoxidation in acetonitrile– aqueous biphasic media has been reported by Page et al. [6].

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In a second review, Rusling [71] emphasizes the green impact of microemulsions and also their ability to direct electrosynthesis toward unexpected products. Despite not using a true emulsion, Chiba et al. [75] have used biphasic electrosynthesis to perform Diels–Alder reactions of electrogenerated quinones in the close proximity of water by employing an SDS surfactant. Under conventional conditions, the quinone would be unstable in water, but it is the effect of the local hydrophobic conditions that allows the reaction to proceed. A similar study by Chiba et al. [76] demonstrated the application of PTFE fibers to form hydrophobic regions for reactions. Often, emulsions are stabilized by surfactant molecules, which can cause complications in the work-up. Alternatively, metastable emulsions can also be generated by power ultrasound [77–79] or via ULTRA-TURRAX [69] emulsification. These emulsions are generated only during the external application of high-shear mixing but return to separate bulk phases when the agitation is removed. This methodology has the benefit of an emulsion during a reaction but is able to be easily separated postreaction in its bulk phase form. Asami et al. [78,79] reported the applications of acoustically generated emulsions in electrosynthesis. The benefits of the emulsion in this case are that the redoxactive substrate is anodically converted in an electrolytic phase, whereas the oxidation-sensitive nucleophile is maintained in a separate organic phase without electrolyte to avoid the competing oxidative pathway (see Figure 8.5 and Table 8.3). Furthermore, the separated electrolytic phase was chosen as ionic liquid, 1-ethyl-3-methylimidazolium tetrafluoroborate (EMIM+BF4−), due to its ability to dissolve the pyrrolidine substrate, remain immiscible with the acetonitrile nucleophilic phase, and not dissolve the nucleophile. In a single-phase system, the allyltrimethylsilane nucleophile would be preferentially oxidized, leading to lower yields. A further benefit of the biphasic system is that the electrolytic phase may be recovered for reuse, and the product remains in the acetonitrile for easy extraction. A related methodology recently developed for biphasic electrosynthesis is the use of the triple-phase boundary electrode–aqueous electrolyte–organic solvent [80]. This technique Nucleophile (without electrolyte)

Nu

Nu

Nu

Acoustic emulsification

P Nu

S Nu S

S

S

Electrolysis

P

S S

Nu S

Nu

Emulsion

S = Substrate N u = Nucleophile S+= Carbocation P = Product

Electrolytic medium

N

S+

Nu

Emulsion

P

Nu

S S

–

+

SiMe3

+

–2e–, –H+ N EMIM BF4

COOMe

COOMe

–2

(1 mmol)

(10 equiv)

0.2 mA cm , Pt – Pt

FIgURE 8.5 Schematic description and reaction scheme for the electrolysis methodology based on in situ emulsification of reagents and mass transport enhancement in 1-ethyl-3-methyl-imidazolium tetrafluoroborate employing a divided cell. (From Asami, R. et al., Chem. Commun., 244, 2008.)

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TAbLE 8.3 Data for Electrolysis Processes based on In Situ Emulsification of Reagents and Mass Transport Enhancement in 1-Ethyl-3-Methyl-Imidazolium Tetrafluoroborate Employing a Divided Cell Entry

Substrate

1

Charge (mol equivalent) 2

N

COOMe

COOMe 2

43 N

N COMe

COMe 3

3

50 N

N

COOMe

COOMe N

yield (%) 70

N 2

4

Product

4

COOMe

N

66

COOMe

Source: Asami, R. et al., Chem. Commun., 244, 2008.

involves placing the electrode at the interface of electrolyte phase and organic solvent phase allowing, for example, reductive electrosynthesis to be completed in an organic phase without the need for intentionally added electrolyte [81,82].

V. ULTRASOUND IN ORgANIC ELECTROSyNTHESIS Ultrasound is defined as a sound wave with a frequency above the range of human hearing, typically about 20 kHz and higher. The frequency and power used often play a key role in the various applications of ultrasound. Generally, a distinction can be made between two frequency bands of ultrasound [83]. For electrosynthesis, frequencies between 20 and 100 kHz (power ultrasound) are generally used due to their larger mechanical, agitation, or cavitation effects [84,85]. Ultrasound is highly beneficial for synthetic electrochemistry by providing a high degree of mass transport, especially for slow diffusers [86], gained through two key mechanisms: (1) acoustic streaming, where a directed turbulent jetflow of bulk material is in effect causing fast material transport to the electrode surface; and (2) cavitation or microstreaming, where the effect of the ultrasonic wave causes a compression and rarefaction cycle in the bulk liquid leading to implosive voids being created at the electrode surface [87]. Upon collapse, these voids lead to a roughening of the electrode surface and additional localized mass transport. Often, ultrasonic bath systems are employed, but these introduce some degree of irreproducibility, and therefore, horn emitters have been introduced to apply power ultrasound more locally to the electrode surface. Three geometries of horn probe sonication in electrochemistry are considered in a review by Compton et al. [7] as the face on, the side on, and sonotrode case (see Figure 8.6). In general, for power ultrasound, the face on configuration is chosen for direct mass transport and

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Electrode

Ultrasonic horn emitter

Organic Electrochemistry

Face on

(a)

Sonotrode

Electrode

Side on Electrode

(b)

(c)

FIgURE 8.6 Three ultrasonic horn emitter geometries for the application of power ultrasound (a) face-on, (b) side-on, and (c) in sonotrode configuration to electrodes for high mass transport. (From Compton, R.G. et al., Electroanalysis, 9, 509, 1997.)

cavitation at the electrode surface. For an effective electrolysis, it is key that the faradaic current is kept as high as possible, and this is generally achieved by using a high degree of mass transport. This high value of faradaic current is directly related to the diffusion layer thickness, and Marken et al. [88] have shown that in an acoustically generated mass transport system, the mass transport–limited diffusion layer thickness is governed by the solvent viscosity. Marken et al. have also demonstrated a simple ultrasound-enhanced reductive synthetic approach for a range of activated alkenes to their corresponding alkanes [89] as emulsions in water without the need for organic solvent. Mass transport enhancement effects are not the only benefit from ultrasound in heterogeneous electrochemical processes. Ultrasound can also induce mechanistic changes of electrochemical processes [90,91] and stop electrode surfaces from being fouled by polymers and other solid materials in a depassivation effect, for example, in Birch electroreduction [92]. The depassivation is mainly achieved through the cavitation mechanism and dependent on frequency, sonication power, electrode geometry, and solvent.

VI.

MICROWAVES IN ORgANIC ELECTROSyNTHESIS

In spite of the popularity of microwave activation in organic synthesis [93], there have been only relatively few attempts to apply microwaves to electrosynthetic reactions. Microwave electrochemical experiments have been carried out by Compton et al. [94] predominantly at microelectrodes. Recent reviews summarize some of the progress [8,95]. For metallic microelectrodes, a focusing effect is observed that allows high-intensity microwave radiation to be applied directly to the dielectric at the tip of the electrode [96]. However, for larger film electrodes, for example, tin-doped indium film electrodes [97], direct heating of the electrode in a flow-through reactor dominates (see Figure 8.7), similar to cases where carbon- or boron-doped diamond electrodes are employed [98]. Radiofrequency heating [99] of platinum film electrodes has also been proposed as a way of enhancing electrochemical processes in flow reactors.

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Application of Ionic Liquids, Emulsions, Sonication, and Microwave Assistance (c)

(e)

341

(d)

Collection vessel

Sample reservoir

(a)

(b) (f ) Microwave field

FIgURE 8.7 Schematic drawing of a microwave-enhanced flow electrolysis cell with (a) the port of the microwave cavity, (b) a Teflon or glass flow cell, (c) the working electrode lead-in inserted into the cavity, (d) the upstream reference electrode, (e) the downstream counterelectrode, and (f) the working electrode in the flow field. (From Rassaei, L. et al., Electrochim. Acta, 54, 6680, 2009.)

VII.

SPECIAL CONDITIONS IN SPECIFIC APPLICATIONS

Apart from activation methods based on heat, ultrasound, and microwave, there have been some reports exploiting new combinations such as plasma processes at the surface of ionic liquid media [100,101]. Plasma discharge processes that are known to allow ozone generation and polymer deposition could be of wider use with an ionic liquid phase to supply reagents or capture products.

REFERENCES 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19.

Anastas, P.; Warner, J. Green Chemistry: Theory and Practice; Oxford University Press: New York, 1998. Little, R. D.; Moeller, K. D. Electrochem. Soc. Interf. 2002, 11(4), 36–42. Rusling, J. F.; Zhou, D. L. J. Electroanal. Chem. 1997, 439, 89–96. Frontana-Uribe, B. A.; Little, R. D.; Ibanez, J. G.; Palma, A.; Vasquez-Medrano, R. Green Chem. 2010, 12, 2099–2119. Paddon, C. A.; Atobe, M.; Fuchigami, T.; He, P.; Watts, P.; Haswell, S. J.; Pritchard, G. J.; Bull, S. D.; Marken, F. J. Appl. Electrochem. 2006, 36, 617–634. Page, P. C. B.; Marken, F.; Williamson, C.; Chan, Y.; Buckley, B. R.; Bethell, D. Adv. Synth. Catal. 2008, 350, 1149–1154. Compton, R. G.; Eklund, J. C.; Marken, F. Electroanalysis 1997, 9, 509–522. Rassaei, L.; Marken, F. Chim. Oggi–Chem. Today 2009, 27, 14–16. MacFarlane, D.; Forsyth, M. The Handbook of Ionic Liquids: Electrochemistry; Wiley-Blackwell: New York, 2011. Wilkes, J. S. Green Chem. 2002, 10, 73–80. Izutsu, K. J. Solid State Electrochem. 2011, 15, 1719–1731. Chu, D. B.; Zhou, Y.; Zhang, X. J.; Li, Y.; Song, Q. Prog. Chem. 2010, 22, 2316–2327. Hallett, J. P.; Welton, T. Chem. Rev. 2011, 111, 3508–3576. Wasserscheid, P.; Welton, T. Ionic Liquids in Synthesis, 2nd edn.; Wiley-VCH: Weinheim, Germany, 2008. Tzschucke, C. C.; Markert, C.; Bannwarth, W.; Roller, S.; Hebel, A.; Haag, R. Angew. Chem., Int. Ed. 2002, 41, 3964–4000. Hapiot, P.; Lagrost, C. Chem. Rev. 2008, 108, 2238–2264. Yoshida, J.; Kataoka, K.; Horcajada, R.; Nagaki, A. Chem. Rev. 2008, 108, 2265–2299. Ab Rani, M. A.; Brant, A.; Crowhurst, L.; Dolan, A.; Lui, M.; Hassan, N. H.; Hallett, J. P.; Hunt, P. A.; Niedermeyer, H.; Perez-Arlandis, J. M.; Schrems, M.; Welton, T.; Wilding, R. Phys. Chem. Chem. Phys. 2011, 13, 16831–16840. Gale, R. J.; Osteryoung, R. A. J. Electrochem. Soc. 1980, 127, 2167–2172.

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20. Wang, H.; Zhao, C.; Bhatt, A. I.; MacFarlane, D. R.; Lu, J. X.; Bond, A. M.; ChemPhysChem 2009, 10, 455–461. 21. Abbott, A. P.; Harris, R. C.; Ryder, K. S.; D’Agostino, C.; Gladden, L. F.; Mantle, M. D. Green Chem. 2011, 13, 82–90. 22. Zhao, C.; Burrell, G.; Torriero, A. A. J.; Separovic, F.; Dunlop, N. F.; MacFarlane, D. R.; Bond, A. M. J. Phys. Chem. B 2008, 112, 6923–6936. 23. Barrosse-Antle, L. E.; Bond, A. M.; Compton, R. G.; O’Mahony, A. M.; Rogers, E. I.; Silvester, D. S. Chem. Asian J. 2010, 5, 202–230. 24. Bornemann, S.; Handy, S. T. Molecules 2011, 16, 5963–5974. 25. Meng, Y.; Aldous, L.; Compton, R. G. J. Phys. Chem. C 2011, 115, 14334–14340. 26. Dong, B.; Xu, J. K.; Zheng, L. Q. Prog. Chem. 2009, 21, 1792–1799. 27. Sekiguchi, K.; Atobe, M.; Fuchigami, T. J. Electroanal. Chem. 2003, 557, 1–7. 28. Martiz, B.; Keyrouz, R.; Gmouh, S.; Vaultier, M.; Jouikov, V. Chem. Commun. 2004, 6, 674–675. 29. Tomilov, A. P.; Turygin, V. V.; Kaabak, L. V. Russ. J. Electrochem. 2007, 43, 1106–1122. 30. Palombi, L. Electrochim. Acta 2011, 56, 7442–7445. 31. Barhdadi, R.; Comminges, C.; Doherty, A. P.; Nédélec, J. Y.; O’Toole, S.; Troupel, M. J. Appl. Electrochem. 2007, 37, 723–728. 32. Gaillon, L.; Bedioui, F. Chem. Commun. 2001, 1458–1459. 33. Herath, A. C.; Becker, J. Y. Electrochim. Acta 2008, 53, 4324–4330. 34. Lagrost, C.; Hapiot, P.; Vaultier, M. Green Chem. 2005, 7, 468–474. 35. Villagrán, C.; Banks, C.; Pitner, W.; Hardacre, C.; Compton, R. G. Ultrason. Sonochem. 2005, 12, 423–428. 36. Chiarotto, I.; Feroci, M.; Orsini, M.; Feeney, M. M. M.; Inesi, A. Adv. Synth. Catal. 2010, 352, 3287–3292. 37. Cruz, H.; Gallardo, I.; Guirado, G. Green Chem. 2011, 13, 2531–2542. 38. Feroci, M.; Chiarotto, I.; Orsini, M.; Sotgiu, G.; Inesi, A. Adv. Synth. Catal. 2008, 350, 1355–1359. 39. Sawamura, T.; Inagi, S.; Fuchigami, T. J. Electrochem. Soc. 2009, 156, E26–E28. 40. Hasegawa, M.; Ishii, H.; Cao, Y.; Fuchigami, T. J. Electrochem. Soc. 2006, 153, D162–D166. 41. Hasegawa, M.; Ishii, H.; Fuchigami, T. Green Chem. 2003, 5, 512–515. 42. Jabbar, M. A.; Shimakoshi, H.; Hisaeda, Y. Chem. Commun. 2007, 1653–1655. 43. Mellah, M.; Gmouh, S.; Vaultier, V.; Louikov, V. Electrochem. Commun. 2003, 5, 591–593. 44. Mellah, M.; Zeitouny, J.; Gmouh, S.; Vaultier, M.; Jouikov, V. Electrochem. Commun. 2005, 7, 869–874. 45. Barhdadi, R.; Courtinard, C.; Nédélec, J. Y.; Troupel, M. Chem. Commun. 2003, 1434–1435. 46. Duran Pachon, L.; Elsevier, C. J.; Rothenberg, G. Adv. Synth. Catal. 2006, 348, 1705–1710. 47. Tang, M. C. Y.; Wong, K. Y.; Chan, T. H. Chem. Commun. 2005, 1345–1347. 48. Ho, K. P.; Wong, K. Y.; Chan, T. H. Tetrahedron 2006, 62, 6650–6658. 49. Buzzeo, M. C.; Klymenko, E. V.; Wadhawan, J. D.; Hardacre, C.; Seddon, K. R.; Compton, R. G. J. Phys. Chem. B 2004, 108, 3947–3954. 50. Yang, H. Z.; Gu, Y. L.; Deng, Y.; Shi, F. Chem. Commun. 2002, 274–275. 51. Lagunas, M. C.; Silvester, D. S.; Aldous, L.; Compton, R. G. Electroanalysis 2006, 18, 2263–2268. 52. Schäfer, H. J. Comptes Rendus Chim. 2011, 14, 745–765. 53. Subramaniam, B. Ind. Eng. Chem. Res. 2010, 49, 10218–10229. 54. Hyde, J. R.; Licence, P.; Carter, D.; Poliakoff, M. Appl. Catal. A: Gen. 2001, 222, 119–131. 55. Giovanelli, D.; Lawrence, N. S.; Compton, R. G. Electroanalysis 2004, 16, 789–810. 56. Fuchigami, T.; Tajima, T. Electrochemistry 2006, 74, 585–589. 57. Grinberg, V. A.; Mazin, V. M. Russ. J. Electrochem. 1998, 34, 223–229. 58. Chanfreau, S.; Cognet, P.; Camy, S.; Condoret, J. S. J. Supercrit. Fluids 2008, 46, 156–162. 59. Matsuda, T.; Harada, T.; Nakamura, K. Green Chem. 2004, 6, 440–444. 60. Anderson, P. E.; Badlani, R. N.; Mayer, J.; Mabrouk, P. A. J. Am. Chem. Soc. 2002, 124, 10284–10285. 61. Hiejima, Y.; Hayashi, M.; Uda, A.; Oya, S.; Kondo, H.; Senboku, H.; Takahashi, K. Phys. Chem. Chem. Phys. 2010, 12, 1953–1957. 62. Zhao, G. Y.; Jiang, T.; Wu, W. Z.; Han, B. X.; Liu, Z. M.; Gao, H. X. J. Phys. Chem. B 2004, 108, 13052–13057. 63. Sullenberger, E. F.; Dressman, S. F.; Michael, A. C. J. Phys. Chem. 1994, 98, 5347–5354. 64. Feess, H.; Wendt, H. Performance of electrolysis with two-phase-electrolyte. In Techniques of ElectroOrganic Synthesis; Part III, Weinberg, N. L.; Tilak, B. V. (eds.), John Wiley & Sons: New York, 1982, pp. 81–178. 65. Mackay, R. A.; Texter, J. Electrochemistry in Colloids and Dispersions; Wiley-VCH: Weinheim, Germany, 1992.

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9

Combinatorial Electrochemistry and Miniaturization Kevin D. Moeller

CONTENTS I. II.

Introduction .......................................................................................................................... 345 Split–Pool-Type Methods...................................................................................................... 347 A. Cation Pools .................................................................................................................. 347 B. Microflow Processing and the Cation-Flow Method .................................................... 353 III. Parallel Synthesis Methods................................................................................................... 355 A. Initial Electrochemical Efforts .....................................................................................355 B. Application to Synthesis: Bulk Methods ...................................................................... 357 C. Microflow Applications to Parallel Synthesis............................................................... 358 D. Microelectrode Array–Based Approaches....................................................................360 IV. Conclusions and a View for the Future................................................................................. 366 References ...................................................................................................................................... 367

I. INTRODUCTION Efforts to discover and develop new molecules that can perform a particular function (either biologic or catalytic) often benefit from the screening of molecular libraries. Use of the libraries enables multiple candidate structures to be rapidly screened, a process that can accelerate both the discovery of initial lead compounds for development and the collection of data that can inform subsequent work. Such efforts depend on the availability of synthetic tools that allow for construction of the libraries. The result has been the development of combinatorial strategies for molecular synthesis. These strategies typically fall into one of two general categories: split–pool methods (Figure 9.1) and parallel synthesis methods (Figure 9.2). Split–pool applications to both biological problems [1] and catalysis development [2] have been recently reviewed. Early examples of parallel synthesis methods [3] and more recent applications [4] have also been reviewed. Both split–pool and parallel synthesis approaches have advantages. The split–pool method is best for building larger libraries because it exponentially increases the number of molecules synthesized with each new step, while parallel synthesis methods lead to much smaller, spatially isolated libraries that can be more readily characterized. With both types of libraries, modern methods for miniaturization have paid significant dividends, especially in the area of microarray technology. Libraries of molecules are only as useful as the methods available to screen their activity. Miniaturization aids these efforts by allowing for automation of the screening process, simultaneous evaluation of multiple molecules that enables higher throughput analysis, lower costs, and the use of biologically relevant amounts of material (e.g., the screening of whole-cell extracts) [5]. As with traditional synthetic strategies, electrochemical methods offer unique opportunities to expand the scope of combinatorial methods. For split–pool-type strategies, electrochemical methods 345

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Organic Electrochemistry

Split

M M M M

M

Reaction 1

M1

M

Reaction 2

M2

M

Reaction 3

M3

M

Reaction 4

M4

Recombine

Pool

Pool with 16 new compounds

Recombine

Split pool (third-iteration)

M1 M2 M3 M4

Split–pool (second-iteration)

M11 M21 M31 M41

Reaction 1

M1 M3

M2 M4

M12 M22 M32 M42

Reaction 2

M1 M3

M2 M4

M13 M23 M33 M43

Reaction 3

M1 M3

M2 M4

M14 M24 M34 M44

Reaction 4

M1 M3

M2 M4

Reactions 1–4 Recombine Pool with 64 new compounds

FIgURE 9.1

Split–pool synthesis.

M

Reaction 1

M1

Reaction 2

M12

M

Reaction 2

M2

Reaction 3

M23

M

Reaction 3

Reaction 4 M

M3

M4

Four pathways lead to four new molecules, etc. Reaction 4

Reaction 1

M34

M41

FIgURE 9.2 Parallel synthesis strategies.

are used to generate highly reactive intermediates in a manner that allows them to be subdivided and then used to trigger a variety of synthetic transformations. Miniaturization of such efforts with the use of microflow reactors imparts unique selectivities to these reactions. For parallel synthesis efforts, multiple electrodes are used to guide the synthesis of library members. The result is the generation of products that are each associated with the electrode used in their generation. In these cases, miniaturization leads to spatially isolated, addressable libraries that can take advantage of microarray technology to both build and monitor the libraries. In this chapter, recent developments in combinatorial electrochemistry are highlighted along with the potential the techniques hold for making significant contributions to the larger field of molecular library synthesis and evaluation.

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Combinatorial Electrochemistry and Miniaturization

II.

SPLIT–POOL-TyPE METHODS

A.

CATION POOLS

Electrochemical strategies aimed at split–pool-type syntheses began with the development of the cation-pool method by Suga, Yoshida, and coworkers. Reviews of both early studies [6a] and more recent efforts [6b] have appeared. The cation-pool method first appeared in an intriguing 1999 report that demonstrated the stability of electrochemically generated N-acyliminium ions in nonnucleophilic solvents at low temperature (Scheme 9.1) [7]. The N-acyliminium ions were generated by the anodic oxidation of an amide in dichloromethane solvent. The reactive intermediate could then be stored as long as the temperature was kept below −50°C. Subsequent treatment of the N-acyliminium ion with an allylsilane, enol ether, aryl, or malonate nucleophile led to the addition products. Two points about the experiment were of particular note. First, the reaction allowed for a one-pot oxidation—nucleophilic addition sequence involving the amide. Hence, the more cumbersome, less sustainable two-step protocol that required the generation and isolation of a methoxylated amide followed by a Lewis- or Bronsted-acid-catalyzed addition to a regenerated N-acyliminium ion could be avoided. In the one-step procedure, the amide oxidation reaction was completed before the addition of the nucleophile. Therefore, the method was still compatible with the use of electronrich nucleophiles with oxidation potentials lower than that of the amide. Second, formation of the stable cation pool allowed for a split–pool approach to combinatorial synthesis. To this end, the N-acyliminium could be subdivided into smaller pools and then each new pool treated with a different nucleophile in order to generate a series of derivatives from the amide starting material. In effect, the cation-pool method allowed for a reactive intermediate to serve as the branch point for a combinatorial synthesis effort. This initial paper triggered a number of applications and advances of the cation-pool approach. In 2000 [8], the work was extended to include the electrochemical generation of oxonium ions (Scheme 9.2, Equation a). The oxonium ions were generated by the oxidative cleavage of a carbon–silicon bond at low temperature. Methods involving the use of electroauxiliaries have been reviewed [9]. The oxonium ions generated decomposed above −50°C, but as in the case of the N-acyliminium ions, at low temperature, they could be subdivided and then trapped with a number of nucleophiles. SiMe3 –2e–

+

–H+

N

N N

CO2Me

CO2Me

CO2Me Stable cation pool 87%

OCOCH3 O

O

O N O

MeO2C O

71% N

N CO2Me 88%

SCHEME 9.1

The use of cation pools.

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MeO2C 72%

348

Organic Electrochemistry + OMe

OMe C8H17

TMS

OMe

–2e–

SiMe3

–Me3Si+

C8H17

C8H17 80% (14 examples, 54–84%)

(a) MgBr –2e–

N

–H+

CO2Me

+N CO2Me

N MeO2C 79% (16 examples, 45–79%)

(b)

SCHEME 9.2 Cation pools of electro-generated oxonium ions (a) and N-acyliminium Ions (b), and the use of reactive nucleophiles.

The initial cation-pool work was also expanded in terms of the nucleophile used. Since the nucleophile used in the reaction does not need to be compatible with the electrolysis, any nucleophile can be utilized. In 2001 [10], Yoshida, Suga, and coworkers demonstrated that N-acyliminium ions generated using the cation-pool method can be trapped with organometallic reagents (Scheme 9.2, Equation b) (see also Chapter 35). Cation pools can also be trapped with alkenes and acetylenes to affect a net cation carbohydroxylation of the π-system (Scheme 9.3) [11]. For these reactions, the carbamate serves as a protecting group for a primary amine. Hence, the reactions enable the synthesis of amino alcohols and amino ketones. Dimethylsulfoxide has also been used as a nucleophile to trap cation pools. In this case, the reactions generate sulfoxonium ions (Scheme 9.4) [12]. Workup with a triaryl amine then affords a chemical oxidation step that generates a ketone product. In an another intriguing application (Scheme 9.5), a cation pool was electrochemically reduced to form radicals [13]. The radicals can be used in order to generate homodimers, or they can be treated with electron-poor olefins in order to generate carbon–carbon bonds (see also Chapter 17). Cation pools have been used to not only generate radical intermediates but also trap radical intermediates (Scheme 9.6) [14]. In this chemistry, hexabutyldistannane serves to propagate a radical-chain process. It is involved in the initial generation of an alkyl radical that then adds to the N-acyliminium ion in the cation pool leading to the formation of a nitrogen radical cation. Ph

Ph SiMe3

N O

O

Anodic oxidation – “SiMe3” Bu4NBF4 CH2Cl2

N O

a)

b) H2O, Et3N

a) R b) H2O, Et3N

R1

N

SCHEME 9.3

R O

Cation pools and carbohydroxylation reactions.

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R1

R1

N O

R O

Yields = (four examples, 60–85%) R = H, Ph R1 = SiMe3, Ph, OAc, C5H11

O

Ph

O

R1

+ O

OH

Ph

R

Yields = (five examples, 47–81%) R = H, Me, Ph R1 = SiMe3, Ph, C6H13

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Combinatorial Electrochemistry and Miniaturization

NHTs

Me + S Me O H

Anodic oxidation (2.1 F), 0°C

O Et3N, 35°C

N

Bu4NBF4 1:2 DMSO/CH2Cl2

1h

Ts

N Ts 86%

8 Cyclizations were performed using sulfonamide, alcohol, aryl, and styrene nucleophiles. Yields ranged from 52–90%

SCHEME 9.4

Cation pools, DMSO, and the formation of ketones. +2e– TfOH

Anodic oxidation –2e–, –H+ N

+N CO2Me

–72°C

CO2Me

CO2Me

N CO2Me

CO2Me (Nine examples, 39–84%)

+e–

SCHEME 9.5

N

N

CO2Me

CO2Me

75%

Electrochemical reduction of cation pools. Anodic oxidation –2e–, –H+ –72°C

N CO2Me

+N CO2Me

Bu3SnSnBu3 R–I –20°C CH2Cl2

Bu3SnSnBu3 –20°C CH2Cl2

N CO2Me

N CO2Me

R N CO2Me

(Three examples, 73–86%)

N CO2Me

81%

SCHEME 9.6

Cation pools and tin-initiated radical chain reactions.

The radical cation then oxidizes the hexabutyldistannane to form a new radical cation that fragments in the presence of the tetrafluoroborate electrolyte to generate fluorotributylstannane and the tributyl tin radical. The tin radical then reacts with the starting alkylhalide to make the alkyl radical and start the process over again. The reactions worked best with alkyliodides and the fivemembered ring N-acyliminium ion. N-Acyliminium-based cation pools have also been used to generate new ring skeletons. In one interesting example, the addition of nucleophiles to the sequentially generated intermediates in a cation pool has been used to generate spirocyclic products (Scheme 9.7) [15]. The chemistry

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Organic Electrochemistry

Anodic oxidation –2e–, –SiMe+3

SiMe3 N

SiMe3

–78°C

CO2Me

SiMe3

+N

Me3Si

n

n N

(or a Grignard reagent)

CO2Me

SiMe3

CO2Me 72–88% Anodic oxidation –2e–, –SiMe3+ –78°C

n N

m CO2Me

Me3Si Ring closing metathesis

n N

(or a vinylzinc reagent)

m

CO2Me

77–92%

m +N CO2Me

n

(3 examples, 46–67%)

SCHEME 9.7 Cation pools and the synthesis of spirocyclic compounds.

SiMe3 Bu

–2e–, –SiMe+3

N

–78°C O

O

a)

Bu +N O

R1 O

Me

Me

SCHEME 9.8

b) H2O

R2 Bu O

R1

N O

R2

(11 examples, 68–88%)

Diels-Alder reactions with pooled cations.

takes advantage of a pair of silyl groups in order to guide the regioselectivity of two successive N-acyliminium ion formations toward the same carbon. In each case, the cation generated is trapped with either an allylsilane or a vinyl zinc to afford a bis-olefin product that is then cyclized with a ring-closing metathesis reaction. The chemistry was used to complete an efficient synthesis of cephalotaxine. In a second example, cation pools have been used to accomplish Diels–Alder reactions that utilize N-acyliminium ions as dienes for an electron-rich dienophile (Scheme 9.8) [16]. The use of the cation pool is essential for the success of these reactions because the electron-rich dienophile is not compatible with the amide oxidation used to generate the N-acyliminium ion. In one intriguing aspect of this work, the reaction yield for cycloadditions using styrene as the dienophile benefited greatly from micromixing (see the following text) the olefin and the N-acyliminium ion. This approach decreased the amount of polymer generated. The polymer is generated from a stepwise addition of the styrene to the N-acyliminium ion leading to a benzylic cation. With poor mixing, this cation can undergo an addition to a second styrene, an event that starts the polymerization. With micromixing, the effective concentration of styrene is lower. This favors the intramolecular reaction that finishes the net cycloaddition reaction. The effects of miniaturization can have a number of additional benefits for the reactions, something that will be discussed in Section II.B. Cation pools can also be used to generate new polymers. For example, cation pools of diphenylcarbenium ions have been used to synthesize dendrimer structures (Scheme 9.9) [17]. The use of a cation pool in this manner provides an opportunity to synthesize dendrimers using a split–pool approach. Another variation in the cation-pool method uses the technique to generate highly reactive intermediates that are then used to remotely trigger reactions (Scheme 9.10). Both the method [18] and its application to oxonium ion formation [19] have been reviewed.

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Combinatorial Electrochemistry and Miniaturization

F

F

Ar

F

F

Anodic

Ar

Ar

Oxidation

Ar

+ Cation pool Ar

Ar

Ar

SiMe3

SiMe3

Ar = p-F-C6H4–

Ar

Anodic Oxidation Ar

etc. SiMe3 Ar

Ar

Ar

Ar

Ar + Ar

Ar

SiMe3

Cation pool

SCHEME 9.9 Cation pools and the synthesis of polymers. Substrate Starting material

R.I.+ I+

Product

R.I. = Reactive Intermediate

SCHEME 9.10

A new approach to generation and use of reactive intermediates.

The result is a method for the generation of reactive intermediates that cannot be made with a direct electrolysis. The method again opens up new opportunities for the use of split–pool approaches to synthesis. In principle, either the initial reactive intermediate generated (R.I.+) or the second intermediate generated (I+) can be used as the branching point for such an effort. One intriguing application of this chemistry is illustrated in Scheme 9.11 [20]. In this example, an anodic oxidation of ArSSAr is used to generate ArS(ArSSAr)+ as a reactive intermediate that in turn triggers the formation of an oxonium ion from an alkoxy thiol acetal. The indirect oxidation was advantageous over the direct oxidation to form the oxonium ion because the direct electrolysis reaction is too slow. Hence, the oxonium cannot be accumulated prior to decomposition. This problem can be solved by first generating and accumulating the more stable ArS(ArSSAr)+, and then using this reactive species to rapidly generate the desired oxonium ion. The oxonium ion can then be employed in a split–pool synthesis to synthesize a number of olefin addition products. The reactions provided a significant improvement over previous efforts that had utilized the direct oxidation of silylated ethers to generate the oxonium ions [19]. A similar mediated approach to oxonium ion formation using the ArS(ArSSAr)+ intermediate has been used to initiate electrochemical glycosylation reactions [21]. OMe Anodic oxidation Bu4NBF4/ ArSSAr CH2Cl2

C8H17

SPh

ArS(ArSSAr)+BF4–

TMS

OMe

OMe C8H17

+

C8H17 (Seven examples, 42–91%)

SCHEME 9.11

An indirect oxidation to form oxonium ions.

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352

Organic Electrochemistry

Anodic oxidation Bu4NBR4/ ArSSAr CH2Cl2

, MeOH ArS(ArSSAr)+BR4–

Me OMe

(R = F/Ar = p-FC6H4)

SAr 68%

Ph

a) R1 (R = C6F5/ Ar = p-MeOC6H5)

R1

R2

b) Et3N ArS SAr (R = F/Ar = p-FC6H4) (Four examples, 66–84%)

SAr ArS

R2

Ph 77% (Three other examples, 58–84%)

SCHEME 9.12

Alkene substrates and the formation of sulfur substituted products. O

Anodic oxidation Bu4NBR4/ ArSSAr CH2Cl2

R1

R2

SAr

O

ArS(ArSSAr)+BF4– R1

R2

R = C 6F5/Ar = p-FC6H5

SAr

(Six examples, 62–88%)

SCHEME 9.13

An indirect approach to triggering cyclization reactions.

In the examples highlighted by the reaction in Scheme 9.11, a split–pool synthesis could be accomplished by taking advantage of the oxonium ion generated. Split–pool syntheses can also be accomplished by subdividing the initial ArS(ArSSAr)+ intermediate (Scheme 9.12) [22,23]. In these reactions, a thiol group from the reactive species is incorporated into the product. A similar approach was also used to generate oxonium ions and initiate cyclization reactions from a series of substrates (Scheme 9.13) [24]. Six examples of the reaction have been accomplished with yields ranging from 62% to 88%. In a related application, an indirect electrolysis reaction has been utilized to oxidize furan rings on solid supports (Scheme 9.14) [24]. The anodic oxidation was used to generate “Br +” that in turn Anodic oxidation nBu4NBr 1:1 MeOH/1,4-dioxane

O N H

O

3

O

O

C-anode, 15 mA/cm 40 F

O N H

2

HO

Indirect oxidation of solid phase substrates.

© 2016 by Taylor & Francis Group, LLC

O

O OMe

Five equiv. LiOH 20/1: 1,4-dioxane/water RT, 2 days

Source of “Br+”

SCHEME 9.14

O

3

OMe O OMe (Six examples, 53–63%)

OMe

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Combinatorial Electrochemistry and Miniaturization

oxidized furans tethered to polystyrene beads. After cleavage of the product from the beads, the reactions afforded the methoxylated furans in good yield and excellent purity.

B.

MICROFLOW PROCESSING AND THE CATION-FLOW METHOD

As illustrated for the Diels–Alder reaction in Scheme 9.8, electrochemical reactions can greatly benefit from miniaturization and the use of a microflow reactor [25,26]. Microflow reactors have channels that are micrometer size in scale, a situation that allows for the downsizing of chemical reactions. Such reactions offer four potential advantages over larger, more typical batch processes. First, they offer the opportunity for micromixing [27], the aspect of the process that improved the earlier Diels–Alder reaction. Micromixing is accomplished by having two solutions flow into a micromixer that is comprised of a collection of narrow channels that are arranged in either a parallel or a series format. This increases the contact area between two solutions while reducing the diffusion length necessary for complete mixing, a scenario that results in more uniform mixing of reactive species. Second, microflow reactors offer improved temperature control. This advantage steps from the greater surface-to-volume ratio of the reactors that enables faster heat exchange between the reactor and the bulk solution. Third, microflow reactors offer better control over reactions that take place on the interface between the catholyte and the anolyte in an electrolysis reaction. This advantage again stems from the greater surfaceto-volume ratio in the microflow reactor that in this case leads to improved phase-boundary conditions. Finally, microflow reactors offer control over the contact time a reactive intermediate has with the electrolysis conditions. For a sensitive intermediate, short contact times can dramatically improve yields. Many of the recent efforts to develop microflow electrolysis reactions began with the realization of how the advantages offered by a microflow reactor might improve the synthetic reactions derived from cation pools (Scheme 9.15) [25]. This work has been reviewed [26]. By taking a stable cation pool and then introducing it into a microflow reactor, a cation flow can be generated. This cation flow can then be treated with nucleophiles, a method that can improve mixing of the cation with the nucleophile, allow for shorter reaction times with sensitive cations, and provide better temperature control over the often exothermic process. The initial examples of this approach used the electrolysis reaction to make a pool of N-acyliminium ions that were subsequently trapped with allylsilane and enol ether nucleophiles [25,26c,28]. In this work, Yoshida and coworkers nicely demonstrated the utility of the method for combinatorial synthesis by utilizing a combination split–pool/parallel synthesis approach. To this end, an iminium ion pool was subdivided and then processed in parallel fashion with several different nucleophiles. The nucleophilic trapping of an N-acyliminium ion in a microflow reactor benefited greatly from the improved mixing associated with the experiment [29]. Consider the Friedel–Crafts

Substrate

Cation pool

Micro mixer

Nucleophile

SCHEME 9.15

Microflow techniques and the use of cation pools.

© 2016 by Taylor & Francis Group, LLC

Product

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Organic Electrochemistry O

O MeO

N

Anodic Oxidation

Bu

MeO

+ Bu N

Bu MeO

Micromixer

OMe

SCHEME 9.16

N

Bu OMe + MeO

OMe N OMe Bu O N

OMe

OMe

MeO

O

OMe

O

SiMe3

OMe

Batch mixing

33%

:

33%

Micromixing

92%

:

4%

OMe

Microflow and the Friedel–Crafts alkylation reaction.

reaction shown in Scheme 9.16. In this example, the alkylation of 1,3,5-trimethoxybenzene with an N-acyliminium ion led to the formation of a mixture of both mono- and disubstituted products. When the reaction was accomplished in a batch reactor, roughly equal amounts of the products were formed. However, when the reaction was conducted in a microflow reactor equipped with a micromixer, both the yield and the selectivity of the reaction improved. Using the microflow channel, a 92% yield of the monoalkylated product was formed along with only 4% of the dialkylated product. The observation that micromixing improved the chemoselectivity of the reaction proved to be general. For example, the use of a micromixer also led to a dramatic improvement in the yield of monosubstituted product generated from the iodination of electron-rich aryl rings [30]. Interestingly, the same reactive N-acyliminium ion intermediate employed in the chemistry illustrated in Scheme 9.16 can be used to trigger polymerization reactions with enol ether monomers [31]. The reactions again benefited from the generation of a cation pool of the substrate and then the fast mixing associated with the microflow reactors. In this case, the combination of cation-pool chemistry and the microflow reaction enabled better control over the monodispersity of the polymer. The better distribution resulted from the ability to carefully generate and control the concentration of the reactive initiating group. It was important to quench the polymerization in a uniform manner so that the polymerization reaction ran for a uniform length of time. This was accomplished by controlling the path length of the reaction in the microflow channel before introduction of a quenching reagent (Scheme 9.17). O MeO

+ Bu N Cation pool Variable path length Micromixer

O

Micromixer Bu

OMe N

N(i-Pr2) n OR

OR

(i-Pr)2NH, CH2Cl2 (quencher)

R = n-Bu, i-Bu, t-Bu

SCHEME 9.17

Microflow and the synthesis of polymers

© 2016 by Taylor & Francis Group, LLC

Cooling bath

OR

355

Combinatorial Electrochemistry and Miniaturization O + Bu N

MeO

Bu N

CO2Me

Micromixer O

Micromixer

N CO2Me

O

82% SiMe3

N CO2Me

SCHEME 9.18

(Other allylsilane, allyl stannane, and ketene acetal olefins were studied)

Microflow and the use of multicomponent reactions.

ArS(ArSSAr)+ –B(C F ) 6 54

OBn BnO BnO

+ O

OBn Micromixer

Micromixer

BnO BnO

OBn BnO BnO

SCHEME 9.19

O

SAr

OBn

Micromixer

ROH

OBn

Et3N

O

OR

OBn

Yield = 67–98% R = Me, sugar

Microflow and the synthesis of glycosides.

In a similar manner, the addition of a second reagent to a microflow reaction affords the opportunity to conduct multicomponent reactions (Scheme 9.18) [32]. In this case, the second coupling component is added downstream of the initial nucleophile added to the N-acyliminium ion. The cation-flow strategy also proved useful for the generation of reactive reagents. For example, the use of [ArS(ArSSAr)]+ has proven particularly useful for the generation and synthetic application of glycosyl cations (Scheme 9.19) [33]. In these experiments, the yield of the product obtained was highly dependent on both the resident time and the temperature of the reaction. Since the reaction benefited from low temperature and shorter retention times, it appeared that the stability of the oxonium ion was an issue. Overall, the approach proved compatible with the generation of a variety of oxonium ions and the trapping of those oxonium ions with alcohol, allylsilane, and enol ether nucleophiles.

III.

PARALLEL SyNTHESIS METHODS

A. INITIAL ELECTROCHEMICAL EFFORTS The reactions outlined earlier fall into the split–pool category because they use a single electrode to generate a reactive intermediate that is then subdivided and used to initiate a number of subsequent reactions. The methods are attractive because they convert a single intermediate into a variety of products. Many of the reactions are also compatible with a parallel synthesis format that uses multiple electrodes to generate a variety of different intermediates that are then individually carried forward. Of course, the techniques can be combined in that the intermediates generated in a parallel process can be used to start split–pool syntheses.

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Organic Electrochemistry

Parallel approaches to electrochemistry have been frequently used in the discovery of new metalbased catalyst systems [34]. In these efforts, electrochemistry is typically used as an analytical technique for rapidly screening the utility of catalysts that are synthesized using traditional means. The pioneering work in this area was reported by Smotkin and Mallouk in 1989 [35]. In this effort, catalysts for methanol oxidation were screened by pipetting various metal salts and aqueous sodium borohydride onto carbon paper. The carbon substrate was conductive but not catalytic for the oxidation of methanol. The result was an array electrode that could be used to screen the catalysts for activity. This was done by stepping up the potential of the array electrode in the presence of a fluorescent indicator that was luminescent in acid. Since methanol oxidation generates acid, the presence of the indicator showed the locations on the array electrode where methanol was oxidized most readily. In this way, all of the catalysts could be screened in a single experiment. A very similar method was used to discover methanol-tolerant nonplatinum electrocatalysts for use in direct methanol fuel cells [36]. More traditional parallel synthesis approaches have also been used for catalyst discovery. For example, a series of Pd-binary catalysts for the reduction of oxygen were spotted onto a working electrode [34b]. The electrode was then covered with a PTFE mask that had 88 holes with a diameter of 5 mm and a depth of 10 mm. The holes were aligned with the catalyst on the working electrode. A Pt-wire counter electrode and Ag/AgCl reference electrode were then placed into one of the holes, and the oxygen reduction reaction activity measured for the catalyst composition at that site. Following the electrochemical evaluation in one of the holes, the counter and reference electrodes were moved to the next site. In this way, a series of new catalysts could be systematically prepared and rapidly evaluated. Speiser and coworkers have used a similar parallel approach to discover new catalysts for conducting synthetic organic reactions [37]. In a recent example, a series of homogeneous ruthenium(II) hydrogenation catalysts were explored by first dissolving them in an electrolyte solution and then placing them in the wells of a microtiter plate [37d]. Electrode bundles were then moved using computer control [37b] and placed into each of the wells one at a time. Cyclic voltammograms were recorded for each catalyst. The voltammograms obtained in this manner could be correlated to the success of the catalysts leading to a high-throughput method for identifying new catalysts. The importance of such efforts is driving the development of a number of interesting new technologies. For example, Crooks and coworkers have reported a new method for screening oxygen reduction catalysts using bipolar electrodes (Scheme 9.20) [38]. To this end, a bipolar indium tin oxide (ITO) electrode is coated with a series of Ag strips on one end and a catalyst for the reduction at the other. The functionalized bipolar electrode is then placed in an electric field that creates an anode on the end with the silver strips and a cathode at the catalytic end. If the catalyst on the cathode can do the desired reduction of oxygen, then a current is passed that reduces the oxygen at the cathode and oxidizes the Ag strips on the anodic side of the bipolar electrode. The potential drop in the bipolar electrode is proportional to distance between the electrodes so the oxidation takes place first at the Ag strip located furthest from the catalyst. When this strip is consumed, the next strip becomes active lowering the overall potential of the system. The more active the catalyst is for

O2 H2O

Ag+

+

– Ag strips

SCHEME 9.20

ITO

Catalyst

Bipolar electrodes and the screening of oxygen reduction catalyses.

© 2016 by Taylor & Francis Group, LLC

357

Combinatorial Electrochemistry and Miniaturization

oxygen reduction, the lower the potential that can be tolerated by the reaction leading to the consumption of more Ag strips. In this way, the activity of the catalyst can be directly ascertained from an examination of the Ag strips consumed. By placing a series of functionalize bipolar electrodes in the electric field, a number of catalysts can be evaluated at the same time. The bipolar electrode approach is also applicable to the synthesis and evaluation of other materials. Inagi and Fuchigami have used bipolar ITO electrodes in order to synthesize gradient-doped conducting polymer films [39]. To this end, the ITO electrodes were coated with a thin film of the conducting polymer and then placed in an electric field. This led to a gradient potential in the bipolar electrode, and hence a gradient potential being applied to the polymer. The gradient potential led to varying levels of dopant being inserted into the polymer matrix with the amount of dopant being dependent on the distance from the ends of the ITO electrode.

B. APPLICATION TO SYNTHESIS: BULK METHODS Parallel electrochemical methods are also useful for the synthesis of small organic molecules. Yudin and coworkers have employed parallel electrosynthesis methods to conduct both oxidation and reduction reactions. The work has been reviewed [40]. As an example, a strategy for the oxidation of amides, carbamates, and sulfonamides is illustrated in Scheme 9.21 [41]. In the setup shown, the electrolyses were conducted in 16 individual reaction vials (or wells in a Teflon block) that were fitted with a Teflon cap with 16 matching openings. In between the Teflon cap and the vials was placed a stainless steel plate with 16 stainless steel electrodes welded in position to match the vials. These stainless steel electrodes served as the counter electrode in each of the reactions. Into each vial was placed a working electrode using the opening in the Teflon cap. In this way, a galvanostatic electrolysis could be conducted in each vial. The reaction setup was used to both make a small library of alkoxy-substituted products and rapidly optimize the reaction conditions for a particular reaction. The same reaction setup was also employed for a reduction reaction (Scheme 9.22) [42]. In this example, diamines were synthesized by the Pd(0)-mediated electroreductive coupling of imines. In order to minimize the cost of the parallel electrolysis setup, a stainless steel cathode was used along with a sacrificial aluminum anode. The sacrificial anode was employed in order to avoid Anode Cathode (stainless steel plate beneath the cap)

SCHEME 9.21

O

Ot-Bu

The parallel synthesis of N-acyliminium ions. Cathode

Al anode NH2 2+

Pd0

R1

R1

Pd

2/3 Al3+

NH2 2 R1

SCHEME 9.22

–2e– ROH

N

NH + TFA H

Parallel synthesis and reduction reactions.

© 2016 by Taylor & Francis Group, LLC

N

OR

Ot-Bu O 80–95%

358

Organic Electrochemistry Series format

– Cathode

+

–

+

–

–

+

Anode

+

n

Electrochemical cells

TMS R Cbz

N H

N O

OMe CH2Ph

R = H, Me, CH2Ph Cbz = PhCH2OC(O)

SCHEME 9.23

RVC anode, Pt cathode 21.0 mA, 2.0 F 0.03 M Bu4NBF4, MeOH

O

Each substrate was oxidized in two separate cells for a total six simultaneous oxidations

MeO R Cbz

N H

O N

O

OMe CH2Ph

Yields = 74–91%

A series approach to N-acyliminium ions.

oxidation of the product amines. The reactions were superior to conventional coupling reactions with activated metals in that they led to significantly less reduction of the imine without coupling. Presumably, this change in the reaction occurred because of the low working potential of the cathode due to the use of the sacrificial anode, a feature that affords careful control of Pd(0) generation. The oxidation of silylated amides using both a parallel electrochemical approach and a seriestype format has also been reported (Scheme 9.23) [43]. The silyl group was added to the amides in order to expand the scope of the substrates that can be oxidized. For example, with the use of a silyl electroauxiliary, a single amino acid in a peptide can be oxidized. The compatibility of the chemistry with combinatorial synthesis was initially demonstrated using a parallel synthesis strategy directly analogous to the method illustrated in Scheme 9.21. This approach allowed for individual control of each electrolysis reaction. However, for these reactions, independent control of the reactions was not necessary. With the electroauxiliary, the oxidation potential of the substrates did not vary greatly, and a simple series format could be used (Scheme 9.23). In this experimental setup, the cathode of one electrolysis reaction is hooked to the anode of the next. A single power supply is used, and the amount of current passed through each of the cells is identical. The series format has also proven compatible with the oxidation of carbamates without the electroauxiliary [44]. In this case, both pyrrolidine- and piperidine-derived carbamates were methoxylated. The reaction took advantage of a solid-supported base [45] to generate the electrolyte for the reaction. This change enabled a high-throughput separation of the products generated.

C.

MICROFLOW APPLICATIONS TO PARALLEL SYNTHESIS

As with the earlier split–pool syntheses, microflow channels offer a number of advantages for parallel reactions. In one powerful development, microflow channels have been used to conduct laminar-flow-based reactions (Scheme 9.24) [46]. This method relies on the channel being small enough to ensure that the flow through it is both stable and laminar. The result is a stable interface between two flows introduced into the microchannel. Substrates introduced in Flow 1 do not reach the anode, and substrates introduced in Flow 2 do not reach the cathode. In this way, cations can be generated at the anode and then allowed to react with nucleophiles at the interface between the flows without exposing the nucleophiles to the oxidation. The nucleophiles can either be introduced

© 2016 by Taylor & Francis Group, LLC

359

Combinatorial Electrochemistry and Miniaturization Flow 1

Cathode Subs.

Nuc–

Subs.

E+

Anod e

Flow 2

SCHEME 9.24

Product (Nuc-E)

A laminar-flow approach to parallel synthesis.

in Flow 1 or generated at the cathode (see the following text). The reaction setup has a major advantage in that it allows for reactions with unstable cations. In the first example of this reaction, an N-acyliminium ion generated at the anode was quenched with an allylsilane nucleophile that was introduced into the reaction via Flow 1 [46,47]. As an alternative, the nucleophile can be added to the product from the electrolysis following the electrochemical reaction [25]. Once again, the laminar flow in the channel controlled the rate of product formation. This second setup allowed more time for the formation of the desired electrophile, an approach that can be beneficial if the electrophile generated is stable. The laminar-flow method has also been used to generate anions (Scheme 9.25) [48]. In the reaction shown, the chemoselective reduction of either an aldehyde or an allylchloride is conducted independent of the reduction potentials of the two groups. This was accomplished by controlling which channel the substrates were introduced into. For example, if the aldehyde was introduced through Inlet 1, then it was reduced to a radical anion that in turn underwent reaction with the allylchloride to form A as the major product. If the allylchloride was introduced through Inlet 1, then it was reduced to an allylanion that added to the aldehyde to predominately form B. The reaction illustrates how laminar flow in the microreactor can be employed to control the chemoselectivity of the reaction. Because of the closeness of the electrodes in a microflow cell, the reactions are ideal for paired syntheses that take advantage of both the anodic and cathodic processes of the electrolysis. A review on this effort has appeared [49]. To this end, Atobe, Fuchigami, and coworkers used the laminarflow-based reaction design highlighted in Scheme 9.25 to add benzyl anions to ketones [50]. In these transformations, the cathode was used to make the anion, while the anode was used for the oxidation of a secondary alcohol. O Flow 1 Method A Ar H Cl

Method B

Cathode Subs.

OH

Nuc– Ar Anode

Flow 2 Cl Method A O

Method B Ar

SCHEME 9.25

H

Laminar-flow and the synthesis of anions.

© 2016 by Taylor & Francis Group, LLC

A. Method A

OH vs. Ar B. Method B

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Organic Electrochemistry

N+ SiMe3

N

CO2Me Anode

CO2Me Micromixer N Cathode

+

Ph

CO2Me

Cl

Ph

70% Ph

TMSCI

SiMe3

SCHEME 9.26 Divided flow reactors and paired electrolyses. Anode

80 μm

O + MeOH

MeO

O

OMe

+ H2

Cathode Max. yield = 98%

SCHEME 9.27 Thin-film reactors and paired electrolyses.

In a related example from the Yoshida group, a divided cell microreactor was used to generate both an N-acyliminium ion (at the anode) and an allylsilane nucleophile (at the cathode) [28]. The flow from each chamber was then combined to effect the net addition reaction (Scheme 9.26). An alternative reaction design that flowed substrates as a thin film between two parallel plate electrodes was forwarded by Fuchigami, Marken, and coworkers (Scheme 9.27) [51]. The reactor was used to remove two electrons from a furan group while at the same time generating two equivalents of methoxide and an equivalent of hydrogen. The combination of anodically and cathodically generated intermediates afforded the desired methoxylated furan product. As in the laminar-flow method, this reaction was intriguing from an environmental perspective because it did not require the addition of an electrolyte. Instead, the combination of intermediates generated at the anode and cathode to form the desired product avoided the buildup of charge at the electrodes. As a related reaction design, Yoshida and coworkers used a microflow channel with a porous membrane divider to accomplish the methoxylation of 4-methoxytoluene and N-methoxycarbonyl pyrrolidine without the use of an electrolyte [52]. A pair of alternative microflow reactor designs has been forwarded by Kristal and coworkers [53]. In these reactors, the electrolysis is conducted by passing the substrate through a parallel array of bipolar electrodes [54] (one example is shown in Scheme 9.28). The result is a series of parallel channels that each comprises a galvanostatic cell. The approach enables a parallel electrochemical method with greatly simplified electric circuitry.

D.

MICROELECTRODE ARRAY–BASED APPROACHES

In principle, microelectrode arrays can also be used to conduct the parallel synthesis of a molecular library. A description of the arrays cited in the following has been published [55]. With a microelectrode array, each electrode in the array is used to facilitate the synthesis of a molecule in the library. The result is unique in that it builds each molecule in the library proximal to a unique, individually addressable microelectrode in an array. The electrodes can then be used to monitor the binding properties of the molecules. For example, consider the electrochemical impedance experiment outlined in Scheme 9.29. A review of such experiments as they apply to analytical sensors has appeared [56], and the experiment works well on microelectrode arrays [57]. In the current context, an array

© 2016 by Taylor & Francis Group, LLC

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Combinatorial Electrochemistry and Miniaturization

Out

–+ –+ –+ –+ –+

+ + + + +

Anode

–+ –+ –+ –+ –+

–+ –+ –+ –+ –+

– – – – –

Cathode

In

SCHEME 9.28

Parallel synthesis using bipolar electrodes. Fe2+ O O

S

M1 Fe3+

O Signal

O

S

Fe2+

M2R

Fe2+

O O

M2

M3 S Fe3+

Anode

SCHEME 9.29

Auxiliary cathode

A microelectrode array approach to molecular signaling.

functionalized with ligands for a receptor (M1, M2, and M3) is submerged in a solution containing an iron redox couple along with a Pt counter electrode. The iron redox couple is then oxidized at the array and reduced at the counter electrode leading to a current that can be measured at each electrode of the array. The targeted receptor (M2R) is then added to the solution above the array. When the receptor binds a molecule on the surface of the array, it blocks the iron redox couple from reaching the anode below. The current drops off at that electrode leading to a signal that can be measured. The result is a method for monitoring the binding of the molecules on the surface of the array and the receptor as the events happen. Of course, the challenge of setting up such an experiment is placing or building the molecules in the molecular library proximal to individual electrodes in the array, especially when the array contains upward of 12,544 electrodes/cm2. One approach to solving this problem has been forwarded by Heller and coworkers [58]. In this work, free-field electrophoresis was used to deliver charged reagents to specific reaction sites on an array. Each site on the array was comprised of an attachment layer, a permeation layer, and an underlying direct current microelectrode. The attachment layer was used to bind a monomer to the site, and the permeation layer was used to prevent reagents

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Organic Electrochemistry

generated at the underlying electrode from reaching the surface. Hence, all reactions at a given site on the chip resulted from the charged reagents delivered to the site. That delivery was accomplished by using the electrode to attract and then concentrate the charged reagent at the site. Once there, the reagent triggered a transformation involving the molecule bound to the attachment layer. For example, reagents were delivered to selected sites on the array in order to carry out the deprotection of a monomer. Subsequent coupling reactions conducted on the array occurred only at sites selected for the deprotection. A similar method for building peptides was reported by Beyer and coworkers [59]. In this work, activated amino acids were embedded into polar particles. Electrical fields patterned by individual pixel electrodes were then used to deliver the particles to various sites on an array. Strong adhesive forces then held the particles in place on the surface even when new electrical fields were applied. In this way, multiple amino acids were delivered to multiple different sites on the array. After placement of the particles on the array, the particles were melted leaving behind the activated amino acid substrate. The ensuing coupling reactions then added new amino acids to the peptides being constructed at each location. The method is advantageous in that the coupling reactions are all done in parallel. This allows for the construction of a molecular library with a minimal number of coupling steps. An alternative strategy is to use the electrodes at selected sites in an array to synthesize reagents and catalysts for doing desired reactions at the site. Since electrodes can be used to synthesize acids, bases, oxidants, reductants, nucleophiles, electrophiles, transition metal catalysts, etc., such an approach would allow for a wide range of synthetic chemistry to be conducted site-selectively on an array. Efforts along these lines began with the site-selective generation of acid and the synthesis of DNA oligomers [60,61]. The chemistry started by coating the array with a sucrose surface [62] that was then functionalized with a DNA primer. The electrodes in the array were then used to site-selectively generate acid by the oxidation of water. Base in solution confined the acid to the electrodes selected. The acid generated led to the removal of a dimethoxytrityl protecting group from the 3′-hydroxy of the DNA, and then the alcohol used to add a nucleotide to the DNA oligomer. Repeating the process allowed for the synthesis of an oligomer. The overall strategy is not restricted to the generation of acid [62,63]. It has been shown to work for the site-selective generation of base [57b,64], Pd(II) [65], Pd(0) [66], Cu(I) [67], Sc(III) [68], and ceric ammonium nitrate [69]. In each case, an electrolysis reaction is used to generate the chemical reagent or catalyst, and in each case, a solution-phase reaction is used to scavenge the reagent or catalyst generated before it can migrate to remote sites on the array. The result is a powerful method for the synthesis of a variety of molecules on an array. The site-selective generation of Cu(I) on an array provides an excellent example of how the electrodes in an array can be used as cathodes (Scheme 9.30) [67a]. All such sequences start with coating the array with a porous reaction layer that allows for the attachment of molecules to the surface of the array. For the chemistry outlined in Scheme 9.30, the array was coated with agarose. A substrate was then placed onto the agarose by each electrode in the array with the use of a basecatalyzed esterification reaction [57b]. The base was generated by the reduction of vitamin B12. The resulting radical anion deprotonated the alcohols on the agarose leading to the formation of an alkoxide that in turn added to an N-hydroxysuccinimide ester. The reaction was confined with the use of excess activated ester in the solution above the array. The activated ester reacts with any base (methoxide) that migrates away from the electrode where it was generated. Following the surface functionalization, a Cu(I)-catalyzed click reaction was used to add a benzotriazepine to each of the electrodes in the array. This reaction was conducted by treating the entire array with copper sulfate, disodium bis(bathophenanthroline)-disulfonate ligand, and tetrabutylammonium bromide electrolyte. The copper sulfate was then reduced to the necessary Cu(I) catalyst by setting the electrodes in the array to a potential of −2.4 relative to the Pt-anode counter electrode. In this case, the click reaction was run at every microelectrode in the array, although it can be confined to selected electrodes with the use of oxygen as a confining agent [67]. Following the click reaction, a second Cu(I) reaction was conducted. In this case, the Cu(I) catalyst was used to effect a

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Combinatorial Electrochemistry and Miniaturization

Vitamin B12, Me4NNO3 MeOH, DMF OH –2.4 V, 600 cycles 0.5 s on, 0.1 s off O O N O N3 O

O O

7:2:1 MeCN/DMF/H2O Bu4NBr, CuSO4, Phen. –2.4 V, 400 cycles 0.5 s on, 0.1 s off wholeboard pattern N3 3

N

O N NN

O

3

O

N O

N Me N

N N

Me

7:2:1 MeCN/DMF/H2O Bu4NBr, CuSO4, Ph3P, –2.4 V, 600 cycles 0.5 s on, 0.1 s off checkerboard pattern I pyrene

O O

N

N

N 3

N O pyrene

SCHEME 9.30

N N

Me

(See color insert.) Site-selective Cu(I)-reactions on a microelectrode array.

Heck-type reaction that added a vinylhalide across the imine in the benzotriazepine substrate. The conditions were nearly identical to those used for the click reaction with two exceptions. First, triphenylphosphine was used as the ligand for the reaction in analogy to the known solution-phase reaction. Second, for this reaction, only selected electrodes (a checkerboard pattern) were used for the reduction along with oxygen as a confining agent. The oxygen reoxidized any Cu(I) catalyst in the solution above the array before it could migrate to remote sites on the array. The image provided in Scheme 9.30 shows the success of this confining strategy. The chemistry highlighted in Scheme 9.30 shows not only the utility of the microelectrodes in the array as cathodes but also the use of Cu(I) click reactions as a method for rapidly synthesizing mass spectrometry cleavable linkers [67b]. The array synthesized was analyzed by TOF-SIMS [70]. Under these conditions, the triazole product from the click reaction fragmented and thereby allowed for the mass of the product to be measured. In this way, characterization data can be obtained for molecules synthesized on an array. The microelectrodes in an array are also useful anodes. Two examples serve as particularly effective illustrations. In the first, the electrodes were used to generate Pd(II) in order to effect the oxidation of the agarose polymer coating the array (Scheme 9.31) [65b]. The oxidant was confined to selected electrodes in the array with the use of ethylvinylether as the confining agent. The ethylvinylether confined the Pd(II) by means of a solution-phase Wacker oxidation that rereduced the Pd(II) back to Pd(0). The result was an opportunity to conduct site-selective reductive amination reactions on an array. In the image shown, a checkerboard pattern of electrodes was used for an initial oxidation that was then used to add an amine with a red fluorescent tag to the array (via a reductive amination). The sequence was then repeated using the opposite checkerboard pattern for the oxidation and an amine with a green fluorescent tag for the reductive amination. In the gray scale image shown, the green spots appear as a darker shade of gray. The lighter gray spots are red. No crossover was observed between the electrodes used for the two patterns.

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Organic Electrochemistry Porous hydroxylated polymer

Electrode on

–e– Pd(OAc)2,

OH

Electrode off

16 min

OH

N

NH2–(Red) O NaCNBH3, MeOH 12 h

N 3 Et4NOTs; (7:1) CH3CN/H2O CH3CH2OCHCH2 Br

N+

O

NaO3S

NH2

NaO3S

SO3Na

SO3– O SO NHNH2 (Red)

OH

Electrode off

H N (Red)

Electrode on

OH

–e– Pd(OAc)2,

N 3 Et4NOTs; (7:1) CH3CN/H2O CH3CH2OCHCH2 16 min

(Green) H N (Red)

Br

H N (Red)

NH2–(Green) NaCNBH3, MeOH 12 h

N (Green) H

SCHEME 9.31

O

(See color insert.) Microelectrode arrays and site-selective oxidation reactions.

Similar approaches have been used to generate a range of oxidants including CAN, DDQ, and Cu(II). But the chemistry is not restricted to the generation of stoichiometric oxidants. In the second example of a microelectrode array being used as an anode, the microelectrodes were employed to site-selectively generate Lewis acids on the array. In the chemistry illustrated in Scheme 9.32, the microelectrodes in the array were used to oxidize a Sc(I) precursor. The resulting Sc(III) Lewis acid was used to catalyze a site-selective Diels–Alder reaction [68]. The chemistry worked beautifully O

O NHBoc

O O N

Sc(OTf )3 Me4NNO3 CH2Cl2 3.5 V, 900 cycles 0.5 s on, 0.1 s off Wholeboard pattern

O

NHBoc

O O N

N Ph S

OH + O HN

pyrene

O

H

O

pyrene

H NN H Bu4NPF6, MeOH +3.0 V, 0.5 s on, 0.1 s off, 900 cycles Checkerboard pattern

n N O

SCHEME 9.32

(See color insert.) The use of “safety-catch” linkers on an array.

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on both 1K and 12K arrays. In each case, the Sc(I) precursor for the reaction was generated in the solution above the array by treating a catalytic amount of Sc(OTf)3 with an excess of a 2-arylbenzothiazole. The benzothiazole reduced the Sc(OTf)3 and then served as a confining agent for the ensuing microelectrode array reaction. For the reaction shown in Scheme 9.32, the dienophile for the Diels–Alder reaction was attached to the array with a safety-catch linker [71]. This linker possessed a masked amine that could be deprotected by the generation of acid at the associated electrode, a scenario that was accomplished nicely by the oxidation of diphenylhydrazine [62,63]. Release of the amine triggered lactam formation and cleavage of the molecule from the surface of the array. The result was the ability to recover the Diels–Alder product from the array, a scenario that allowed for the determination of both the stereochemistry and purity of the product. The use of the safety-catch linker strategy is particularly attractive in that it uses the same electrodes employed to monitor the biological behavior of the molecules in the library to recover them from the array. In this way, there is no chance for a loss of fidelity on the array between the signaling and characterization studies. Reactions have also been developed for the placement of more complex biological ligands onto a microelectrode array. Initially, this chemistry focused on the use of a thiol-based Michael reaction (Scheme 9.33) [57b]. The chemistry was confined by controlling the reaction that placed the Michael acceptor onto the array. The image provided illustrates the level of confinement that can be obtained on a 12K array with this strategy. The chemistry proved compatible with the placement of peptides onto the arrays and was used to provide an initial probe of the analytical capabilities of the arrays [57b]. However, the reactions were plagued by the reversibility of the conjugate addition. At a pH 7, molecules placed on an array using a thiol-based Michael reaction migrate from one site to another. The retro-Michael reaction can be stopped at pH 4, but most biological studies require neutral pH. A different approach was needed. To this end, site-selective Cu(I) reactions have proven ideal (Scheme 9.34) [67]. In the example shown, a fluorescent ligand tied to biotin was placed onto an array having 12,544 microelectrodes/ cm2 [57c]. A second set of electrodes was then functionalized with just the linker. Both of the coupling reactions involved a Cu(I)-mediated addition of an amine nucleophile to an aryl bromide surface coating the array [67b]. The Boc-protecting group on the initial substrate was removed during the electrolysis by acid generated at the counter electrode. The Cu(I) catalyst was generated in a fashion identical to that discussed earlier in connection with Scheme 9.28. The rate of the addition was dependent on the nature of the ligand used. As with solution-phase Cu(I)-coupling reactions, the chemoselectivity of array-based coupling reactions with competing heteroatomic nucleophiles is dependent on the nature of the ligand used for copper [67b]. O

O N

O

HS

O

O OH Vitamin B12, Me4NNO3 MeOH, DMF –2.4 V, 300 cycles 0.5 s on, 0.1 s off Single electrode O O

SCHEME 9.33

O

O O

Vitamin B12, MeNNO3 MeOH, DMF –2.4 V, 300 cycles 0.5 s on, 0.1 s off

O S

O

(See color insert.) Functionalizing electrodes with conjugate addition reactions.

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366

Organic Electrochemistry O NH H

HN H

S Diblock copolymer [72]

O

H N Br

On

Ph

4

N

O

R

Boc

CuSO4, Bu4NBr, Ph3P, CH3CN/DMF/H2O

O

N

OFF

H

N H

–1.7 V, 180 s Br

Br

Off

On

H N

O

R1

Boc

HN

CuSO4, Bu4NBr, Ph3P, CH3CN/DMF/H2O, –1.7 V, 180 s

NH H

H S Ph

4

O

N

O

O

N H Me

N H

O

Ph N

O

O

N H

SCHEME 9.34

N H

A biotin functionalized microelectrode array.

The array synthesized with the chemistry illustrated in Scheme 9.34 was used to confirm the utility of the arrays for monitoring molecular interactions in “real time” [57]. In these experiments, the array was submerged in solutions containing an Fe2+/Fe3+ redox couple and various concentrations of streptavidin. For each concentration of streptavidin, a cyclic voltammogram for the iron was obtained. As the concentration of streptavidin increased, the current associated with the iron decreased at electrodes that were functionalized with the biotin. No decrease occurred at either unfunctionalized sites on the array or sites functionalized with the linker but no biotin. The array clearly detected the biotin–streptavidin binding event, a result that demonstrated the potential for the approach as a method for interrogating biological receptors. Similar analytical experiments using peptides that contain arginine–glycine–aspartic acid residues and the integrin receptors they target have proven equally successful [72].

IV. CONCLUSIONS AND A VIEW FOR THE FUTURE Electrochemistry is a versatile tool for organic synthesis that provides a variety of opportunities to solve synthetic challenges in unique ways. This statement appears especially true as the larger field of organic chemistry pushes further into the biological arena and the synthetic challenges encountered

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expand. Such efforts often require the rapid assembly and analysis of molecular libraries. The libraries are assembled by using split–pool and parallel synthesis techniques. For both, electrochemistry offers enticing opportunities. From the use of cation-flow methods that allow for the use of highly reactive intermediates as branching points for combinatorial syntheses to the use of microelectrode arrays that allow for the synthesis of addressable libraries, electrochemistry provides opportunities to rapidly construct molecular libraries in ways that are simply not possible using more traditional methods. From a medicinal chemistry standpoint, this is very important. The availability of unique molecular libraries and unique methods for monitoring them gives rise to the potential for unique, proprietary lead compounds and new opportunities for drug discovery. So unlike previous efforts to apply electrochemistry to the synthesis of natural products, the success of combinatorial electrochemistry will not depend on its adoption by a reluctant organic synthesis community, but rather the opportunities it offers for the application of organic chemistry to the solution of biological problems.

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10

Relations between Microand Macrophenomena Christian Amatore

CONTENTS I. II.

Introduction .......................................................................................................................... 371 Basic Principles .................................................................................................................... 372 A. Homogeneous Chemical Equivalent of Electrolysis Rates ........................................... 372 B. Coulometry: Apparent Number of Electrons Consumed..............................................375 C. Separation of Waves and Selectivity in Electrolysis ..................................................... 379 III. Competitive Reaction Schemes: Optimization of Preparative-Scale Electrolysis................ 381 A. Slow Reaction Schemes ................................................................................................ 381 B. Elementary Method for Fast Competitive Kinetics ...................................................... 383 1. Presentation of the Method .................................................................................... 383 2. Predictive Value of the Method ............................................................................. 386 3. Rationalization of a Series of Results .................................................................... 388 IV. Conclusion ............................................................................................................................ 390 Acknowledgments.......................................................................................................................... 391 References ...................................................................................................................................... 391 There is no art as difficult as the art of observation: a rational intelligent mind together with educated experience, only attainable through practice, are indispensable: not the observer who only visually perceives the object, but rather the observer who sees the different parts and how these parts are connected to the whole. Justus von Liebig

I. INTRODUCTION In Chapter 1, our presentation of electrochemical concepts was focused mainly on an analytical/ kinetic point of view, namely, on the specific aspects related to the electrode and to the solution in its close vicinity. Indeed, as soon as kinetics is considered, the bulk solution may often be regarded as a concentration buffer whose only role is to maintain constant or known reactant concentrations at the external diffusion layer boundary. Obviously, this is not the case when preparative electrochemistry is considered. Indeed, the goal is then to change the solution composition. Yet this composition change is effected through the relay of the diffusion layer, which “transmits” the modifications achieved at the electrode to the bulk solution. Since chemical reactions may take place within the diffusion layer, as discussed in Chapter 1, the way the bulk solution is affected depends largely on the microphenomena occurring in the thin solution layer (5–20 µm, usually) adjacent to the electrode [1]. This results in an extensive coupling between micro- and macrophenomena at electrodes, which is one of the reasons why kinetics, the study of microelectrolysis phenomena, is so developed in electrochemistry.

371

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Organic Electrochemistry

Yet these important relations should not conceal the large differences between micro- and macroelectrolysis. These differences, which were discussed briefly in Chapter 1, have their main origin in the necessity of passing a current through the solution between the working and auxiliary (or counter) electrodes. Obviously, the ensuing problems are quite different when one considers currents in the microampere (or even femtoampere for ultramicroelectrodes) range or in the ampere range as in macroscale laboratory electrolysis.* Other changes are also introduced that originate from the differences in dilution. Under microelectrolytic conditions, very dilute (usually 0.1−10 mM) solutions are considered. Thus, activities and diffusion coefficients, for example, are close to their values at infinite dilution. This is not the case in macroelectrolysis since rather concentrated solutions are used. This may introduce large changes in mass transfer rates as well as in chemical reactivity, and also in such specific factors as migration [3]. These changes in concentration, as well as the larger current densities and electrolysis times, may also result in an increased importance of adsorption, coatings, or pollution phenomena at the electrodes. This is particularly true in aqueous media in which adsorption or chemisorption may play a decisive role in the distribution and selectivity of the products as a function of the electrode material. As understood from this discussion, it is generally unrealistic to believe that microelectrolytic results transpose exactly to macroelectrolytic conditions.† However, provided one keeps these difficulties in mind, microelectrolysis results and approaches constitute extremely useful guides in tackling with macroscale electrolysis. Amazingly, data obtained at ultramicroelectrodes (i.e., of micrometric or smaller dimensions) should transpose with the least difficulty since they may be obtained with concentration or current density (a few amperes per square centimeter) conditions matching those found in preparative-scale electrolysis. Moreover, they allow meaningful results to be obtained in the presence of strong migrational contributions [4,5].

II.

bASIC PRINCIPLES

A.

HOMOGENEOUS CHEMICAL EQUIVALENT OF ELECTROLYSIS RATES

As explained in Chapter 1, the bulk concentrations, C, are affected by the phenomena occurring within the diffusion layer (which is identified, in this chapter, with the stagnant layer imposed by hydrodynamics) through the general relationship dC DA  ∂C   dC  =− +   dt V  ∂x  x =δ  dt chem

(10.1)

where D is the diffusion coefficient A is the working electrode surface area V is the bulk solution volume This equation derives from the conservation of mass fluxes at the diffusion layer/bulk solution interface (compare Equations 1.155 and 1.156 and Figure 10.1). In Equation 10.1, dC/dt is the overall rate of the concentration variations considered, whereas (dC/dt)chem is that arising from possible consumption or production of the species within the bulk solution. The electrochemical rate of production or consumption is given by the first term on the right-hand side of Equation 10.1. It depends on the concentration gradient of the species at the end of the diffusion layer, that is, at x = δ, as well * Note that in the following, we do not discuss the ultraspecific aspects of industrial-scale electrolysis, which generally cannot be dissociated from a particular process [2]. † But this is not particular to electrochemistry, since most of the aforementioned difficulties (with the exception of those related to the current) are encountered when transferring any physical organic result to preparative conditions.

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373

Relations between Micro- and Macrophenomena

i or E

CR(x)

x 0

δ C CR(x)

Electrode

Bulk solution Stagnant layer

δ

FIgURE 10.1 Schematic representation of the electrochemical cell. First insert: concentration profile of the reactant in the stagnant layer in the vicinity of the electrode. Second insert: mass transfer black box at the boundary between the stagnant layer and the bulk solution (x = δ).

as on the ratio of the electrode surface area to the solution volume, A/V. The larger the absolute value of this rate, the smaller the duration of the electrolysis, which is the reason that generally large electrode surface areas and high stirring rates (i.e., small stagnant layer thickness) are used for electrolysis. Note that in Equation 10.1, a single space coordinate is considered, x, the distance normal to the electrode. This necessarily implies that all electrode locations behave identically. In practice, this may be approximated by using symmetrical electrode geometries, with the auxiliary electrode (or separator in divided cells) parallel and of nearly identical size to that of the working electrode. This results in a concentration field as uniform as possible at the electrode surface(s). Note that this field is conveniently visualized when thinking in terms of the electrostatic field that would establish if the electrodes were those of a large capacitor, since both fields obey analogous laws [6]. However, in practice, ideal electrochemical geometries are difficultly compatible with other experimental constraints, so that the potential and current distribution at electrode surfaces may not be uniform. This is also true for the thickness of the stagnant layer since a uniform fluid velocity is difficult to achieve at the electrode surface. For these reasons, the first term on the right-hand side of Equation 10.1 is merely to be considered as an average over the electrode surface; this is also true for the electrode potential E or the electrode current density i/A. Nevertheless, to maintain some generality and simplicity in the following presentation, we assume that the concentration profiles depend only on the distance x from the electrode surface. The problem is then to relate the concentration gradient (∂C/∂x)x = δ at the stagnant layer boundary to that (∂C/∂x)x = 0 at the electrode surface (compare Figure 10.1). This is done via the integration of the differential equations, such as Equation 1.175, which control the concentration profiles within the stagnant layer. In the most simple situation, the species whose bulk concentration is given in Equation 10.1 undergoes no chemical reaction within the stagnant layer. Then from Equation 1.175, its concentration flux is constant within the layer, and thus (∂C/∂x)x = δ = (∂C/∂x)x = 0 = (C−Cx = 0)/δ, where Cx = 0 is the concentration of the species at the electrode surface. Thus, Equation 10.1 is transformed into dC DA  dC  (C − C x = 0 ) +  =−  dt Vδ  dt chem

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(10.2)

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where the constant factor DA/Vδ has the dimensions of a first-order rate constant (s−1). Most of the time, one is interested in exhaustive electrolysis and thus imposes an electrode potential so that Cx = 0 = 0, which in most cases simplifies Equation 10.2 for the reactant: dC DA  dC  =− C +  dt Vδ  dt chem

(10.3)

Interestingly, under such circumstances, the electrochemical consumption of the species is equivalent to a pseudo-first-order chemical reaction taking place in the bulk solution, with a rate constant kelec [7]: kelec =

DA Vδ

(10.4)

Note that it may be surprising that a chemical term is considered in Equation 10.3, whereas we have supposed that no chemical reaction term was involved within the stagnant layer. As explained in Chapter 1, this stems from the fact that the effect of any chemical reaction of rate vchem within the diffusion layer depends on the relative magnitude of vchemδ2/D versus unity. For usual laboratory conditions, δ ≈ 10−3 cm and D ≈ 5 × 10−6 cm2 s−1 and then δ2/D is of the order of 0.2 s. Thus, provided vchem is less than approximately 1 s–1, it has no tangible effect on the concentration profiles of the species inside the stagnant layer, whereas it has a definite effect in the bulk solution due to the long reaction times (usually longer or comparable to half an hour). In practice, it is important to decide when the simplification in Equation 10.2 or 10.3 applies to a given experimental situation. The discussion just presented affords a simple answer to the problem. Indeed, consider the electron transfer reaction in Equation 10.5, possibly followed by a chemical step in Equation 10.6: R + ne ⇌ P ( E 0 )

(10.5)

P → P′ (vchem = k[ P ])

(10.6)

The simplification in Equations 10.2 and 10.3 is equivalent to saying that vchem is much smaller than the mass transfer rate D/δ2, which is of the order of 5 s–1. From Chapter 1, this implies that the R/P redox couple gives a chemically reversible cyclic voltammogram at a scan rate corresponding to the same mass transfer rate. From a dimensionless analysis and Table 1.5, this corresponds to a scan rate v such as Fv/RT ≈ D/δ2, that is, v of the order 0.1 V s–1. Thus, observation for identical or close experimental conditions of a chemically reversible cyclic voltammogram at v ≈ 0.1 V s–1 for the redox couple of interest is sufficient to prove the validity of the assumptions leading to the homogeneous equivalent rate laws in: d[ R ] = −kelec [ R ] dt

(10.7)

d[ P ]  d[ P ]  = + kelec [ R ] +  = + kelec [ R ] − k[ P ]  dt  dt chem

(10.8)

with kelec given in Equation 10.4 and [R] and [P] being the concentrations of R and P in the bulk solution at time t. In the converse situation, that is, when a totally irreversible cyclic voltammogram is observed, Equation 10.7 still applies for R, which is not chemically affected in the sequence considered, but

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375

Relations between Micro- and Macrophenomena kelec

Rbulk

Bulk

Pbulk

etc.

Stagnant layer

Electrode

SCHEME 10.1

Rx = 0 + ne

Px = 0

Macroscopic vs. microscopic description of an electrochemical reaction.

Equation 10.8 does not reflect the variations in the bulk concentration of product P [8]. Indeed, the observation of a chemically irreversible cyclic voltammogram shows that P is too short-lived to have any possibility of escaping that stagnant layer. However, because of mass conservation within the stagnant layer, the outgoing flux of P′, the product of chemical reaction of P in Equation 10.6, is identical to the incoming flux of R. Equation 10.8 is then replaced by [P] ≈ 0, and the following equation results, which is tantamount to the steady-state approximation in homogeneous chemistry: d[ P′] = + kelec [ R ] dt

(10.9)

From this analysis and presentation, it is seen that when single chemical pathways are followed, that is, when no competitive route interferes in the reaction of intermediate(s) within the diffusion layer, the electrolysis affects the different bulk concentrations in a way similar to a first-order chemical reaction, as outlined in Scheme 10.1. This analogy affords an extremely simple way to handle the various problems that may arise during electrolysis. Indeed, they then become exactly identical to those that would be observed for homogeneous chemistry, that is, as if the R/P redox reaction were performed using a homogeneous redox reagent in excess.

B.

COULOMETRY: APPARENT NUMBER OF ELECTRONS CONSUMED

In Section II.A, we have presented the analogy between electrolysis and homogeneous chemistry in terms of an equivalent rate constant kelec = DA/Vδ. Yet it may be surprising that this rate constant does not include a term related to the current. However, the current corresponds to a number of moles consumed per unit of time, that is, to an instantaneous rate, and therefore not to a rate constant. From the description outlined in Scheme 10.1 and Equation 1.141, the current exchanged at the electrode at any time is given in  ∂[ R ]   ∂[ P ]  i = nFAD  = −nFAD     ∂x  x = 0  ∂x  x = 0

(10.10)

When R undergoes no significant chemical reaction within the diffusion layer, its concentration profile is linear, as established previously, and Equation 10.10 is rewritten as i=

nFAD ([R]bulk − [R]x =0 ) δ

(10.11)

Again, when one considers the most frequent situation in which the electrode potential is such as [R]x = 0 = 0, the following is finally obtained: i=

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nFAD [ R ]bulk δ

(10.12)

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Organic Electrochemistry

Comparison to the definition of kelec (Equation 10.4) affords the following relationship between the instantaneous current and the equivalent rate of electrolysis [7]: i = nFVkelec [ R ]bulk

(10.13)

Note that Equations 10.12 and 10.13 are independent of the eventual chemical fate of R or P in the bulk solution but suppose only that R undergoes no chemical reaction within the stagnant layer. The overall quantity of electricity Qt consumed after a duration t of electrolysis is given by the integration versus time of Equation 10.13, which affords t

t

∫

∫

0

0

Qt = i dt = nFVkelec [ R ]bulk dt

(10.14)

From Equation 10.14, [R]bulk is obtained as a function of time, which follows directly from Equation 10.3: d[ R ]bulk  d[ R ]bulk  = −kelec [ R ]bulk +   dt  dt chem

(10.15)

When the last term in Equation 10.15 is negligible, R is consumed only at the electrode. Thus, the charge consumption is easily obtained by considering that each mole of R converted consumes nF coulombs, that is, Qt is given as follows, where C0 is the initial bulk concentration of R: Qt = nFV (C 0 − [ R ]bulk )

(10.16)

On the other hand, from Equation 10.15, [R]bulk decays exponentially with time as [ R ]bulk = C 0e − kelec t

(10.17)

At the end of electrolysis, that is, for t ≫ 1/kelec, one obtains the overall charge consumption, Q, in Equation 10.18. In practice, Equation 10.18 allows the experimental determination of n, the number of electrons consumed per molecule of R during the electrolysis, as in Equation 10.19: Q = nFVC 0 n=

Q C 0 FV

(10.18) (10.19)

But when experimentally possible, one does not proceed from Equation 10.19, because when the R concentration becomes small, that is, near the end of electrolysis, residual currents arising from other electroactive species may contribute significantly to the current and then alter the charge measured. It is then preferable to rely on the plot of Q versus [R]bulk, according to Equation 10.16. Thus, a linear plot should be obtained as in Figure 10.2a, in which extrapolation at [R]bulk = 0 allows a more accurate value of Q to be determined. Note that the ensuing value of n should be identical to that determined from transient or steady-state methods.

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[R]bulk/C 0

Relations between Micro- and Macrophenomena

1.0

1.0

0.5

0.5

II.

I.

Q

0

0

Q

Qth

Charge consumed (Coulombs) (b)

(a)

FIgURE 10.2 Variations in the bulk concentration of the reactant as a function of the charge consumed: (a) in the absence or (b) in the presence of chemical reaction(s) (I) consuming or (II) regenerating the electrolyzed species (initial concentration of the reactant: C 0).

However, it happens frequently that the two figures do not agree, which is a proof that the chemical term in Equation 10.15 is not negligible under preparative-scale electrolysis conditions. Indeed, when the latter is positive, that is, corresponds to a production of R in the bulk solution, as by a regenerating sequence as R + ne → P stagnant layer

(10.20)

P + Z → R + ⋯ bulk solution the instant R concentration is larger than predicted by Equation 10.17. From Equation 10.14, it is seen that Q is then larger than determined via Equation 10.18, which results in a larger number of electrons consumed nap during the electrolysis, when compared with the net electron exchange n at the electrode identified in microelectrolysis: nap =

Q FVC 0

(10.21)

In fact, then nap corresponds not only to the electrolysis of R but also to the formal electrolysis of Z in Equation 10.20. Conversely, when R is consumed in the bulk solution, Q and therefore nap are smaller than predicted in Equations 10.18 and 10.19. A typical example of such behavior is given by the reduction of pyrylium cations (Equations 10.22 and 10.23). Under steady-state or moderate-scan cyclic voltammetry, their reduction involves the net consumption of one electron per molecule, which corresponds to the fast radical dimerization in Equation 10.23. Ph

Ph

(10.22) O

Ph

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+e

O

Ph

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Organic Electrochemistry Ph

Ph

Ph

2 O

O

O

Ph

Ph

(10.23)

Ph

The corresponding rate constant kdim, close to the diffusion limit, has been determined in acetonitrile using ultrafast cyclic voltammetry in the low million volts per second [9a]. However, bulk electrolysis and coulometries indicate an overall consumption of two out of three electrons. This was shown to correspond to the involvement of the slow hydride transfer sequence outlined in the following reactions [10]: Ph

Ph

Ph

O

+

O

Ph

Ph

Slow

O

Ph

Ph

+

O

O

+

Ph

Ph

Ph

Ph

Ph

Ph

Ph

(10.24)

O

Ph

Ph

+ O + B–

O

Ph

O + BH

O

Ph

Ph

(10.25)

Ph

Reaction (10.24) is too slow to be detected in cyclic voltammetry, yet it takes place during the electrolysis to afford finally a 4H-pyran and a bipyranylidene as summarized in the following balanced equation: Ph

3 O

Ph

+ 2 e + B–

Ph

Ph

O

O

Ph

Ph

Ph

+

O

+ BH

(10.26)

Ph

Thus, three pyrylium molecules are consumed for only two electrons exchanged, resulting in nap = 2/3. In this discussion, we have considered the involvement of a slow follow-up reaction of the reactant as a cause for the difference between n and nap. Yet this is not the only situation possible. In most cases, such disagreements are due to the involvement of an ECE kind of sequence such as that featured in Equations 10.27 through 10.29 [11]:

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R + ne → P (E 0 )

(10.27)

D  P → R′  slow versus 2  δ  

(10.28)

R′ + n′e → P′ (E 0′ )

(10.29)

Relations between Micro- and Macrophenomena

379

as well as to their homogeneous analogs in which the heterogeneous electron transfer is replaced by a homogeneous electron transfer (disproportionation sequences). When the interposed chemical step or the disproportionation is slow enough to make no contribution within the diffusion layer, the apparent number of electrons observed in microelectrolysis is n, that is, reflects only the redox process in Equation 10.27. However, the chemical step may be sufficiently rapid when compared with the electrolysis duration to convert quantitatively the primary product P into R′. Thus, provided that n(E 0′ – E 0) > 0, R′ is coelectrolyzed, and the apparent number of electrons exchanged is nap = n + n′. The absolute value of nap may then be larger than that of n when both have the same sign, or smaller when of opposite signs. Similarly, slow disproportionation reactions pulled by a fast protonation step, for example, may affect the nap value, as outlined in the following sequence: R + ne → P 2P ⇌ R + P′ P′ + Z → ⋯ These examples illustrate some of the difficulties that may be encountered when comparing the numbers of electrons exchanged in micro- or macroscale electrolysis. In a few instances, these chemical complications may be detected by important curvatures of the Q versus [R]bulk plots as sketched in Figure 10.2b. Such curvatures are observed when the degree of involvement of the follow-up chemical steps increases during electrolysis. Indeed, because of concentration effects, the effective rate of the chemical step may be slow at the beginning of the electrolysis, and thus a slope corresponding to nap = n is observed. However, when the pertinent concentrations of intermediates increase, the rate of their chemical reaction may become sufficient so that the intermediates reach steady-state behavior, which results in a tendency to reach a slope corresponding to nap = n + n′ while electrolysis progresses. Yet this supposes that the degree of participation changes during the electrolysis, which obviously is not a general fact. Thus, the absence of curvature in the Q versus [R]bulk plots cannot be taken as evidence of no chemical complication, that is, of the fact that n = nap. This is important to remember when using coulometric results, that is, nap values, in the rationalization of electrochemical mechanisms.

C.

SEPARATION OF WAVES AND SELECTIVITY IN ELECTROLYSIS

It is often alleged that the fine-tuning of electrode potential allows high selectivity in electrolysis. This is certainly true for a wide variety of electrochemical situations but may turn out totally inexact in others, as is shown in the following discussion. Indeed, let us consider a system involving two successive reduction waves* at E1 and E2 for a compound involving two reduction sites A and R: A–R. Let us assume, in addition, that electrolysis of related molecules, but without the R center, is known to occur at potentials lying in the same range where the first wave is observed. Thus, an evident rationalization of the observed waves consists in ascribing them to the successive reductions in Equations 10.30 and 10.31: −

A − R + e → i A − R − ⋯ → B − R (k1 ) first wave −

B − R + e → B − R i − ⋯ → B − P (k2 )second wave

(10.30) (10.31)

When the lifetime of the anion radical formed upon the electron transfer in step (10.30) is extremely short, the mechanism in Equations 10.30 and 10.31 is strictly valid, and the electrolysis product obtained at E = E1 is (B – R). Yet when this is not the case, a combination of (A – P) and (B – R) *

Transposition between reductions and oxidations is obvious when discussing on formal grounds, so that the discussion hereafter is developed only for a cathodic process for the sake of simplification of its presentation.

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Organic Electrochemistry

products may be obtained although the electrolysis is performed at the first wave [12,13]. Indeed, − an alternative slow route is now offered to the ( iA − R ) anion radical formed in Equation 10.30 during the long duration of an electrolysis. This slow path involves a possible intra- or intermolecular − electron transfer to the R site to afford the (A − R i ) anion radical: −i

(10.32)

−

A − R ⇌ A − R i− … → A − P −

If the latter is not highly unstable, a back electron transfer occurs to regenerate the ( •A − R ) anion − radical. In the converse situation, (A − R • ) may evolve directly to the A – P product, which corresponds to the overall reduction of the R moiety although the first electron transfer has occurred at the A group. The observed phenomenon is thus tantamount to indirect electrolysis [14], where the A group plays the role of an intramolecular electron mediator for the reduction of the R group. Although not explicitly recognized as such, this is the case for the reduction of aromatic halides, for example, since the initial electron is transferred to the aromatic ring and then internally transferred to the carbon–halogen antibonding orbital, resulting finally in the overall reduction of the carbon– halogen bond. In such situations, the exact nature of the− (B – R) or (A – P) product obtained depends on the − relative chemical stabilities of the A • and R • moieties (rate constants k1 and k2 in Equations 10.30 and 10.31), as well as on the difference between the standard reduction potentials of the two groups, as explained later. Indeed, in a general case, one must consider that the mechanism occurring at the first reduction wave is not as in Equation 10.30 but may involve the sequence in Equations 10.33 through 10.35: −

(10.33)

A−R+e → •A−R k1

A

B

R

(10.34)

A

P

(10.35)

R K k2

A

R

When the homogeneous electron transfer is sufficiently fast to be equilibrated (with an equilibrium − constant K), the formation rate constant of A – P product is k2K when related to • A − R: [B − R]yield k = 1 [A − P]yield k2 K

(10.36)

[B − R]yield k1 (k2 + kb ) = [A − P]yield k2 kf

(10.37)

Thus, the relative yields of the (B – R) and (A – P) products are given in Equation 10.36 and, it is noteworthy, are independent of the electrode potential. In a more general case, the forward kf and backward k b rate constants of the homogeneous electron transfer must be considered, and the relative yields are given in Equation 10.37* and are again independent of the electrode potential, provided it is kept on the first wave. *

Note that in Equation 10.37, the forward and backward rate constants kf and k b of the equilibrium may include a concentration term when the homogeneous electron transfer is bimolecular as in Equation 10.38: •−

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−

A − R + A – R ⇌ A – R + A − R•

(10.38)

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Relations between Micro- and Macrophenomena

III.

COMPETITIVE REACTION SCHEMES: OPTIMIZATION OF PREPARATIVE-SCALE ELECTROLySIS

A.

SLOW REACTION SCHEMES

In electrochemistry, as in other chemical methods, situations in which a single product is obtained are seldom encountered. In most experimental cases, the target product is often accompanied by unwanted side products formed along competitive pathways. Under such conditions, optimization of the target product yield is obviously desirable. Frequently, this optimization is sought on the basis of chemical intuition and of “blind tests” of supposedly critical parameters. Yet is difficult to be certain (except when a 100% yield is obtained!) that the best yield obtained is the maximum one. In the following, we want to show that semiquantitative predictions can be made on the basis of very simple considerations. The goal of these approaches is not to replace the necessary tests, but to provide useful guidelines for the selection of the parameters that affect the yield of the target product. From the preceding presentation, it is easily inferred that different behaviors are obtained when the key species, that is, that at which the different chemical routes are branched, exists only within the stagnant layer or exists in the bulk solution, as schematized in Figure 10.3. In the latter case, the competition problem is analogous to the homogeneous chemical situation, owing to the equivalences established previously. For example, let us consider a reaction sequence like that in Figure 10.3 (case a or b): A+e

B

(E 0)

(10.39)

B

C

(k1)

(10.40)

P1

(k2)

(10.41)

P2

(E 0΄ > E 0)

(10.42)

C +e

For a potentiostatic electrolysis on the plateau of the A/B reduction wave, the reaction sequence in Equations 10.39 through 10.42 is equivalent to the homogeneous mechanism in Equations 10.43 through 10.46, where kelec = DA/Vδ: A

B

(kelec)

(10.43)

B

C

(k1)

(10.44)

P1

(k2)

(10.45)

P2

(kelec)

(10.46)

C

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Organic Electrochemistry

A+e

B

C

A

A

B

B

C

(a) A

B

A

C

0

C δ

0

δ

0

δ

(b) A

B

A

C

(c) A

A

B C 0

X δ

X

(d)

FIgURE 10.3 Schematic representation of the localization of the branching point for an ECC competitive sequence and resulting concentration profiles for (—) A, (- - -) B, and (…) C. (a, b) Slow production of the intermediate key species C; (c) fast production of the key species; (d) reaction in the adsorbed phase.

Thus, the yield of species P1 vis-à-vis P2 is simply expressed as [P1 ]yield k2 k2 = = [P1 ]yield + [P2 ]yield k2 + kelec k2 + DA /Vδ

(10.47)

From Equation 10.47, it is easily seen that maximizing the yield of P1 amounts to decreasing the ratio kelec/k2, that is, of DA/Vδk2. Thus, it obviously follows that one must decrease the electrode surface area and the stirring rate (i.e., increase δ), increase the bulk solution volume, and, when possible, increase k2. The opposite changes are needed when the target compound is P2. Similarly, when B is the compound sought, one wants kelec ≫ k2, that is, a large surface area for the electrode, small bulk solution volume, and high stirring rates (small δ). Thus, owing to the equivalent homogeneous formulation, there are no specific particularities associated with optimization in electrochemistry, as soon as the branching point takes place only

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Relations between Micro- and Macrophenomena

383

within the bulk solution. As explained earlier, this is generally easy to decide on the basis of cyclic voltammograms of the system of interest. Conversely, when the key species exists only in a thin reaction layer adjacent to the electrode (case c in Figure 10.3), the degree of competition between the various routes is greatly influenced by the local concentration profiles of the different species involved. In such cases, general theories have been established that allow a complete and quantitative description of the optimization procedure. However, in the following, we want to present a very simple approach to the problem.*

B.

ELEMENTARY METHOD FOR FAST COMPETITIVE KINETICS

1. Presentation of the Method All the results presented in the following discussion derive from the simple consideration† that when a species is involved in a branched kinetic scheme, the major pathway is the fastest, that is, the main product is that formed according to the pathway corresponding to the highest rate of consumption of the key species. On the other hand, when the key species is formed at a given rate, the faster its consumption, the smaller its concentration. Taking into account these two obvious statements leads to the following conclusion: when two or more pathways compete for the consumption of a given species, the major pathway is that leading to the smallest concentration of this species, when each pathway is considered alone under identical experimental conditions. When two competing mechanisms are considered, a competition parameter may thus be defined as the ratio of the average concentration of the key species obtained when each of these mechanisms is considered alone. Let us denote by C1 and C2 the corresponding concentrations for mechanisms 1 and 2 in Scheme 10.2, when they do not interfere. The competition parameter p = C1/C2 is then such that p ≪ 1 favors the occurrence of mechanism 1 as the main pathway since this condition corresponds to C1 ≪ C2, whereas mechanism 2 is observed when p ≫ 1. In a general case, the analytical dependence of both C1 and C2 on the experimental conditions (concentrations of the reactants and of the coreactants, stirring rate, and so on) and intrinsic factors (e.g., rate constants and diffusion coefficients) may be determined, thus leading to the expression of p as a function of these factors. The effects of these parameters on the overall product distribution are then readily given with the direction and magnitude included in the p expression. Table 10.1 gives the average concentrations to be used for the various intermediates involved in the most basic schemes of electrochemical mechanisms,‡ which may allow the simple derivation of the competition parameters p for most of the possible experimental situations. In the preceding discussion, we have considered for simplicity that only two pathways were possible for the key species. Thus, a single parameter p is defined. However, when more routes are branched for the same species, the situation is a priori more difficult to handle. For three competitive routes, for example, Table 10.1 affords three concentrations C1, C2 , and C3, and thus three parameters p1 = C1/C2 , p 2 = C1/C3, and p 3 = C2/C3 are defined, which are all affected by the experimental and intrinsic parameters. Thus, it is a priori difficult to appreciate the effect resulting of the variation of one of the experimental parameters on the overall competition. Yet it is first noticed that the consideration of only two parameters suffices to describe the competition since they are all given by ratios or products of the two others, as p3 = p 2/p1. *

† ‡

Note that case d in Figure 10.3, which relates to surface reactions between adsorbed and chemisorbed species, is not considered here because of the too specific nature of the electrode–species interactions involved. Yet they are based and supported by a rigorous physicomathematical analysis. See, for example, References 8a,15–21. Note that each of the reactions presented in Table 10.1 may consist of true elemental reactions or of a kinetically equivalent series of elemental steps.

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Organic Electrochemistry etc. Mechanism 1 C Mechanism 2 etc.

SCHEME 10.2

Branching between two mechanisms at the level of a single intermediate.

TAbLE 10.1 “Average Concentrations” of the Different Intermediates Involved in the Main Electrochemical Mechanismsa,b Mechanismb EC ECE

DISP

ECC

Dim 1 Dim 2 (ECE)

Dim 2 (DISP)

Dim 3 (DISP)

ECDim

a b c d

Scheme

Average Concentrationc

A + ne⇌B k B  →C A + ne⇌B k B  →C

[B] = C 0/(kθ)1/2 [C] = C 0 [B] = C 0/(kθ)1/2 [C] = C 0/(kθ)1/2

C + ne → D A + ne⇌B k B  →C kd B + C  →A + D A + ne⇌B k1 B  →C k2 C  →D A + ne⇌B k 2B  →C A + ne⇌B k B + A  →C C + ne → D A + ne⇌B k B + A  →C kd B + C → A + D A + ne⇌B B + A⇌C [K]d kd B + C  →A + D A + ne⇌B k1 B  →C k2 2C  →D

[D] = C 0 [B] = C 0/(2kθ)1/2 [C] = k/kd [D] = C 0 [B] = C 0/(k1θ)1/2 [C] = (C 0/k2)(k1/θ)1/2 [D] = C 0 [B] = C 02/3(3/4kθ)1/3 [C] = C 0/2 [B] = C 02/3/(kθ)1/3 [C] = C 02/3/(kθ)1/3 [D] = C 0/2 [B] = C 02/3/(2kθ)1/3 [C] = (kC 0)2/3/(kdθ1/3) [D] = C 0/2 [B] = (3C 0/4kKθ)1/3 [C] = C 0K[B] [D] = C 0/2 [B] = C 0/(k1θ)1/2 [C] = (C 0/2k2)1/2(k1/θ)1/4 [D] = C 0/2

The concentrations are given for an electrolysis on the plateau of the A/B wave. C0, bulk concentration of A; θ = δ2/D. Usual names, or taken from the nomenclature developed in Savéant’s group. These concentrations are valid provided each resulting figure is lower than approximately C0/10 (except for the final product). K is the equilibrium constant of the reaction considered always equilibrated.

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Relations between Micro- and Macrophenomena δ C0

log p1

CZ

0

2

lo

g

CY

p

3

0 A

C

3

1 B log p2 0

FIgURE 10.4 Kinetic zone diagram and corresponding compass card for a competition between three hypothetical mechanisms, as discussed in the text.

Thus, a convenient presentation of the competition problem consists in the construction of a diagram in which log p1 is the ordinate axis and log p 2 is the abscissa axis, as in Figure 10.4. In such a diagram, kinetic zones corresponding to the predominance of one of the three possible mechanisms may be defined in a similar way as thermodynamic stability zones corresponding to the predominance of a species are defined, for example, as a function of the solution potential and pH (Chapter 1). Thus, mechanism 1, which corresponds to p1 ≪ 1 and p 2 ≪ 1, is located in the lower left corner of Figure 10.4, that is, log p1 < 0 and log p 2 < 0. Mechanism 2 is observed for p1 ≪ 1 and p3 = p2/p1 ≪ 1, and its kinetic zone is then such that log p1 > 0 and log p1 > log p 2. Similarly, the kinetic zone corresponding to the predominance of mechanism 3 is such that p 3 ≫ 1 and p 2 ≫ 1, that is, log p 2 > 0 and log p 2 > log p1 Then three boundary lines delimiting the zone of predominance of each mechanism are constructed in Figure 10.4. Near one of these boundaries, the two delimited mechanisms compete without interference from the third. The general competition, that is, that involving the three mechanisms, is observed near the intersection of the three boundary lines in the central region of Figure 10.4. When such a diagram has been constructed, the effect of each experimental parameter is easily predicted by construction of a compass card, such as that shown in the upper right corner of Figure 10.4. For example, let us assume that p 1 ∝ C 0 δ 2/C y and p 2 ∝ C 0 Cz /δ. Thus, a multiplication of δ by 10, for example, results in an increase of 2 in log p1 and a decrease by 1 in log p 2. A vector corresponding to the “δ effect” can then be constructed in the diagram, which has a projection −1 in the abscissa and +2 in the ordinate. Similarly, other vectors can be drawn for C 0, C y, and Cz to constitute the compass card shown in the upper right corner of Figure 10.4. Such a compass provides an easy visualization of the effect of any variation in the experimental parameters. For example, it is seen in Figure 10.4 that when the system is at point A, an increase in δ results in a trend to shift from mechanism 3 to 2, whereas the same variation results in a tendency to pass from mechanism 3 to 1 when starting from location B or from mechanism 1 to 2 when starting from point C. Such different effects induced by the same experimental variation would have been puzzling and difficult to rationalize, whereas they are easily understood from the kinetic zone diagram and its associated compass card. Similarly, a compass card showing the effect of each intrinsic parameter may be constructed.

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In the following discussion, we want to show how this simple method may be used in practical situations, either to rationalize a series of results or to decide a priori the worthiness of an experimental strategy. To maintain some homogeneity in the presentation, the two aspects are discussed on the basis of examples pertaining to the electrochemistry of aryl halides. 2. Predictive Value of the Method This aspect is developed by addressing the following question: “Is there any chance that biaryls may be formed, via aryl radical dimerization, during the reductive electrolysis of aryl halides?” − The nonconcerted reduction of aryl halides is known [22] to afford a frangible anion radical ArX • (Equation 10.48), which yields a σ-aryl radical by cleaving off the halide ion in Equation 10.49: −

ArX + e → ArX •

(E 0 )

(10.48)

(k1 )

(10.49)

−

ArX • → Ar • + X −

Although these σ-aryl radicals are prone to undergo a facile reduction at the electrode (Equation 10.50) or in solution as in Equation 10.51, one may think of the possibility of impeding these normal pathways to favor their radical dimerization: Ar • +e → Ar −

( E 0′ > E 0 )

−

(10.50)

Ar • + ArX • → Ar − + ArX (kdif )

(10.51)

Ar − + BH → ArH + B−

(10.52)

2Ar • → Ar − Ar (kdim )

(10.53)

Indeed, the dimerization of these σ radicals is expected to be extremely fast and possibly close to diffusion control. Then, intuitively, owing to its bimolecular nature, a concentration effect should favor this step vis-à-vis the reduction at the electrode. Similarly, one could envision that increasing the rate of the anion radical cleavage in Equation 10.49 should decrease its concentration with a concomitant increase in the radical concentration, both effects being in favor of the duplicating step.* The average concentrations for the key species Ar • to be used in the following analysis are obtained from Table 10.1, where the role of Ar • is symbolized by C in the three competing sequences: ECE (Equations 10.48 through 10.50 and 10.52), DISP (Equations 10.48, 10.49, 10.51, and 10.52), and ECDim (Equations 10.48, 10.49, and 10.53). This allows the three competition parameters to be obtained as in Table 10.2, p1 = (ECDim/ECE), p2 = (ECDim/DISP), and p3 = (p1/p2) = (DISP/ECE), from the ratios of the respective average concentrations. Following this procedure, a kinetic zone diagram and a compass card are then constructed in Figure 10.5a. Such a diagram shows that these − considerations on the effect of C0 or of the cleavage rate constant k1 of ArX • are valid only when the point representing the system is in the ECE zone (C0 effect) or in the DISP zone (k1 effect). When starting from an ECDim zone, however, an increase in C0 results in a tendency to shift toward the DISP zone, and an increase in k1 favors reduction via an ECE mechanism. Both effects are counterintuitive but can be understood. For example, when k1 increases, Ar • is produced closer and closer to the electrode surface and then has less and less possibility to dimerize before reaching the electrode to be reduced. As for C0, the effect of the mass transfer rate (i.e., of θ = δ2/D in the compass card) was almost impossible to predict on an intuitive basis. *

Note that Ar• is prone to undergo other facile reaction paths (see Section III.B.3). The concurrent routes in Equations 10.50 through 10.52 then constitute the minimal competitive sequence to be considered.

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Relations between Micro- and Macrophenomena

TAbLE 10.2 Competition Parameters for an ECDim/ECE/DISP Competing Sequence Average Concentration of Ar• Mechanism ECDim

a

Competition Parameter

• b,c

Sequence

[Ar ] 0

1/ 2

Competition Parameter •

ECDim/ECE

1/ 4

(C / 2kdim ) (k1 /θ) C0/(k1θ)1/2

ECDim/DISP

DISP

k1/kd

DISP/ECE

b

c d

•

p1 = (Ar )ECDim/(Ar )ECE

ECE

a

Major Pathwayd

p1 = k13/ 4 θ1/ 4 / (2kdimC 0 )1/ 2 p2 = (Ar•)ECDim/(Ar•)DISP p2 = kd (C 0 )1/ 2 /k13/ 4 (2kdim )1/ 2 θ1/ 4 p3 = (Ar•)DISP/(Ar•)ECE p3 = p1/p2

pi ≪ 1

pi ≫ 1

ECDim

ECE

ECDim

DISP

DISP

ECE

Taken from Table 10.1. k1, rate constant of the aromatic halide anion cleavage; kdim, dimerization rate constant of the σ-aryl radicals; kd, homogeneous electron transfer rate constant (usually kd = kdif, the diffusion-controlled rate constant). C 0, ArX bulk concentration; θ = δ2/D. Both pathways are observed when the value of pi is close to unity (i.e., when ~0.1 ≤ pi ≤ 10).

kd

log p2

log p2

C0 θ

kdim DISP 0

k1

DISP 0 ECE

ECE ECDim

ECDim

0 (a)

0

log p1

log p1

(b)

FIgURE 10.5 (a) Kinetic zone diagram for an ECDim/ECE/DISP competitive sequence (see text, Equations 10.48 through 10.53, and Table 10.2 for definitions of p1, p2, and p3; note that in (b), kd = kdif ). The hatched zone in (b) corresponds to the region without experimental validity.

To illustrate another useful aspect of the method, let us discuss a further experimental importance related to the ECDim zone in Figure 10.5a. Indeed, the representation in Figure 10.5a supposes that any values are possible for p1 and p2. Yet in an actual case, this is not true since the product p1p2 = kdif/2kdim is necessarily larger than 1/2 because kdim cannot exceed the diffusion limit kdif. This implies that the hatched zone in Figure 10.5b, which corresponds to p1p2 < 0.5, has no experimental validity. Thus, no system can be located in this zone regardless of the values of the experimental (C 0, θ, and δ) or intrinsic parameters (D, kdif, kdim, or k1); this shows that almost all the ECDim zone has no experimental significance. Although this method does not allow the maximum yield of biaryl to be determined, it leads to the conclusion that it is necessarily poor and requires extremely precise conditions to be achieved. A complete and accurate analysis [18] of the problem establishes that this maximum yield is of the order of 10–20% at best, which is in complete agreement with maximum experimental yields lying in the range of 7%. However, this does not mean that biaryls cannot be obtained via aryl halide reduction but only that this strategy is not experimentally valid as soon as Ar • is more easily reducible than the parent halide. Yet other approaches to the synthetic problem may be found. An obvious one consists in suppressing the facile reductions of Ar • by using an electrophore as the leaving group X−, which is

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then also the site of the initial electron uptake (Equation 10.48). Ar i may not then be reducible at the reduction potential of the parent pseudohalide, which is shown in benzylic series [23]: Ar2CHS

NO–2

Ar2CHS

NO2 + e

Ar2CH

Dimers

(10.54)

Another approach consists, for example, in using an ArX or a coreactant Ar’X, prone to react with Ar • via the Ullman kind of reaction [24,25]: (10.55)

Ar • + ArX → [ Ar − ArX]• → …

Naturally, each of these possible approaches, as well as many others that may be conceived, may be tested on paper following this method to decide its possible validity and the best experimental conditions to favor it. For example, application of this on paper strategy shows that only moderate yields are obtained for the process in Equation 10.56 when directly reducing the aryl halide at the electrode, whereas interesting yields (i.e., from 80% to 70%) are possible by using a redox mediator for performing the initial ArX reduction [25c]: Ar Ar +

O–

O– H

–e –H+

Ar

O–

(10.56)

3. Rationalization of a Series of Results Figure 10.6 presents the variations in the yield in monodeuterated arene obtained when an aromatic halide is reduced in acetonitrile in the presence of 10% D2O [26]: 10% D2 O ArX + 2e ACN,  → y ArD + (1 − y)ArH

(10.57)

The overall stoichiometry of the reaction shown in Equation 10.57 corresponds to the formation of ArD and ArH via two competitive pathways. ArD is obtained along the sequence presented in Equations 10.48 through 10.51, the resulting σ-aromatic anion being deuterated by D2O, as Ar − + D2O → ArD + DO −

(10.58)

ArH is formed via a H-atom transfer from the ACN solvent to the aryl radical formed upon the mechanism in Equations 10.48 and 10.49, as outlined in the following sequence [26]*: Ar • + CH 3CN → ArH + • CH 2CN (kH ) CH 2CN(+e or ArX • ) → −CH 2CN

(10.60)

CH 2CN + D2O ⇌ CH 2 DCN + DO −

(10.61)

•

−

−

(10.59)

The data in Figure 10.6 show that for the 1-halonaphthalene or 4-halobenzonitrile series, a continuous variation is observed, that is, the yield in ArD increases in the series Cl < Br < I, in agreement *

The scrambling in Equation 10.61 has been shown to play a minimal role during the electrolysis on the basis of experimental isotopic studies [26].

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Relations between Micro- and Macrophenomena

% ArD 100 x

x

CN

50

x

20 10

20

5

2

20

2 2

0

Br

Cl

I

FIgURE 10.6 Experimental variations in the yield of monodeuterated arene ArD as obtained by reduction of the corresponding halide in acetonitrile in the presence of 10% D2O and 0.1 M LiClO4. The number on the curves is the initial concentration (mM) of the halide. (Data from M’Halla, F. et al., J. Am. Chem. Soc., 102, 4120, 1980.)

(at least for the monotonic variation) with usual chemical expectations. But for the 9-anthracene series, a nonuniform variation is observed as a function of the halide. Moreover, the ArD yields obtained for the 9-chloroanthracene are extremely dependent on the concentration or the stirring rate, whereas the effects of these factors are minimal for the bromo derivative and totally negligible for the iodo derivative or for the 1-halonaphthalene or 4-halobenzonitrile series. Such a priori puzzling experimental results may, however, be easily rationalized through the preceding method. From Table 10.1, the two parameters corresponding to the competition between ECE (reactions 10.48 through 10.50 and 10.58) or DISP (reactions 10.48, 10.49, 10.51, and 10.58) sequences and the H-atom transfer ECC sequence (reactions 10.48, 10.49, and 10.59 through 10.61) are obtained as follows*:

*

p1 =

k1 kH

ECC/ECE

(10.62)

p2 =

C 0 kdif θ1/ 2 kH k11/ 2

ECC/DISP

(10.63)

p3 =

p1 p2

DISP/ECE

(10.64)

Note that for the H-atom transfer pathway, steps (10.60) and (10.61), which occur after the branching point Equation 10.59, play no role in the ArD yield. Thus, the sequence of reactions (10.48), (10.49), and (10.59) through (10.61) behaves like an ECC sequence in regard to the competition between the H• versus (+e, +D2O) routes for Ar• reduction.

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Organic Electrochemistry

log p2

kd, C 0

% ArD 100

DISP (ArD) 0

kH

Cl X

Br I

θ

k1 20 mM

ECE (ArD)

50

10 mM 5 mM

X ECC (ArH)

CN

2 mM

X 0

log (k1/kH) (a)

0

log p1

(b)

–12

–8

–4

0

4

8

FIgURE 10.7 (a) Kinetic zone diagram for the ECC/ECE/DISP competitive sequence in the ArD versus ArH formation under the experimental conditions of Figure 10.6 (see text and Equations 10.48 through 10.51 and 10.58 through 10.61; note that for the example discussed, kd = kdif ). (b) Theoretical variations [16] in the yield of ArD as a function of the concentration of ArX (numbers on the curves) and of the rate constants ratio k1/k H. (Experimental data in (a) from M’Halla, F. et al., J. Am. Chem. Soc., 102, 4120, 1980.)

This allows the construction of the kinetic zone diagram in Figure 10.7a according to the previously explained procedure. The location of the experimental systems considered in Figure 10.6 indicates that the benzonitrile or naphthalene derivatives undergo an ECE/ECC competition without interference from the DISP route. Thus, as indicated by the compass card in Figure 10.7a or by the formulation of p1, in Equation 10.62, C0 or θ = δ2/D has no effect on the overall ArD yield, in agreement with the experimental observations. Moreover, for a given aromatic moiety, k H is constant and does not depend on the halide, whereas k1 increases in the order Cl < Br < I. Thus, p1 =  k1/k H increases monotonically in the series, which explains the uniform variations in Figure 10.6. The 9-iodo anthracene also undergoes an ECC versus ECE competition, whereas an ECC/DISP competition is observed for the 9-chloro derivative. From the compass card in Figure 10.7a, it is then seen that C0 or θ considerably affects the ArD yield for the 9-chloroanthracene, whereas there is no effect of these parameters on the iodo derivative. For the bromide, an intermediate situation is observed but is more shifted toward predominance of an ECC than for the iodide or chloride cases, which explains the minimum observed in Figure 10.6 as well as the small variations in the ArD yield with concentration or stirring rate changes [26]. It is thus seen that the simple approach presented here is sufficient to rationalize qualitatively all the a priori puzzling observations in Figure 10.6. But a correct and thorough analysis of the problem on a physicomathematical basis [16] may be achieved and leads to a more quantitative description of these facts. The result of such an analysis is shown in Figure 10.7b, under the form of predicted variations in the ArD yield, for the systems in Figure 10.6. Although more quantitative information is then obtained, it is seen that all the trends observed in Figure 10.7b have been qualitatively predicted via the aforementioned approach, which requires no sophisticated derivations and can be handled “on the back of an envelope.”

IV. CONCLUSION This presentation was designed to discuss the relationships between micro- and macroscale electrolysis. Besides the intrinsic factors arising from scaling-up problems (larger concentrations, larger current densities, and longer reaction times), we have tried to show how the results of microscale electrolysis could be used with large profit, either in devising adequate experimental conditions or in the rationalization of product distributions. However, it is important to decide in such cases if the branching point leading to the different products takes place in the bulk solution or occurs in the

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391

very close vicinity of the electrode. Indeed, in the first case, the problem must be examined in terms of its homogeneous equivalent, the electrolytic reactions being replaced by first-order reactions. Conversely, in the second case, an electrochemical analysis taking into account the concentration profiles in the vicinity of the electrode must be performed to solve quantitatively the distribution problem. However, in such cases, and when only semiquantitative predictions are needed, a very simple approach based upon competition parameters and kinetic zone diagrams proves to be considerably useful.*

ACKNOWLEDgMENTS This work was supported by ENS, CNRS, and UPMC (UMR 8640 PASTEUR). The author wishes to acknowledge Prof. Irina Svir and Dr. Alexander Oleinick for useful comments and suggestions.

REFERENCES 1. For a modern and more correct notion of the so-called Nernst-layer phenomenon see: (a) Amatore, C., Szunerits, S., Thouin, L., Warkocz, J.-S. J. Electroanal. Chem. 2001, 500, 62–70; (b) Amatore, C.; Klymenko, O. V., Svir, I. Anal. Chem. 2012, 84, 2792–2798. 2. See, e.g., (a) Jansson, R. AIChE Symp. Ser. 1983, 79(229), 92–99. (b) Jansson, R. AIChE Symp. Ser. 1983, 79(229), 119–125. (c) Jansson, R. In Electrochemical Cell Design; White, R. E., ed.; Plenum Press: New York, 1984; pp. 175–195. 3. Newman, J. S. Electrochemical Systems; Prentice Hall: Englewood Cliffs, NJ, 1973; pp. 239–253. 4. Wightman, R. M.; Wipf, D. O. In Electroanalytical Chemistry, Vol. 15; Bard, A. J., ed.; Marcel Dekker: New York, 1989. (b) In Ultramicroelectrodes; Fleischmann, M.; Pons, S.; Rolison, D.; Schmidt, P. P., eds.; Datatech Systems: Morgantown, NC, 1987. (c) Wightman, R. M. Science 1988, 240, 415–420. 5. (a) Amatore, C.; Deakin, M. R.; Wightman, R. M. J. Electroanal. Chem. 1987, 225, 49–63. (b) Montenegro, M. I.; Pletcher, D. J. Electroanal. Chem. 1988, 248, 229–232. 6. See, e.g., (a) Reference 3, pp. 340–352; (b) Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Fundamentals & Application, 2nd edn.; John Wiley & Sons: New York, 2001; pp. 421–423. (c) Booman, G. L.; Holbrook, W. B. Anal. Chem. 1963, 35, 1793–1809. (d) Booman, G. L.; Holbrook, W. B. Anal. Chem. 1965, 37, 795–802. 7. (a) Bard, A. J.; Santhanam, K. S. V. In Electroanalytical Chemistry, Vol. 4; Bard, A. J., ed.; Marcel Dekker: New York, 1970, p. 220. (b) Lingane, J. J. J. Am. Chem. Soc. 1945, 67, 1916–1922. 8. (a) Amatore, C.; Savéant, J.-M. J. Electroanal. Chem. 1981, 123, 189–201. (b) Wendt, H. Angew. Chem. 1982, 21, 256–270. (c) Wendt, H.; Plzak, V. J. Electroanal. Chem. 1983, 154, 13–28. (d) Plzak, V.; Wendt, H. J. Electroanal. Chem. 1983, 154, 29–43. (e) Plzak, V.; Wendt, H. J. Electroanal. Chem. 1984, 180, 185–204. (f) Savéant, J.-M. J. Electroanal. Chem. 1987, 236, 31–42. 9. (a) Amatore, C.; Jutand, A.; Pflüger, F. J. Electroanal. Chem. 1987, 218, 361–365. (b) Fabre, C.; Fugnitto, R.; Strzelecka, H. C. R. Acad. Sci. Paris Ser. C 1976, 282, 175–177. (c) Pragst, F.; Ziebig, R.; Seydewitz, U.; Driesel, G. Electrochim. Acta 1980, 25, 341–352. 10. Amatore, C; Jutand, A.; Pflüger, F.; Jallabert, C.; Strzelecka, H.; Veber, M. Tetrahedron Lett. 1989, 30, 1383–1386. 11. See, e.g., (a) Parker, V. D.; Nyberg, K.; Eberson, L. J. Electroanal. Chem. 1969, 22, 150–152. (b) Parker, V. D.; Eberson, L. J. Chem. Soc. Chem. Commun. 1969, 340. (c) Phelps, J; Bard, A. J. J. Electroanal. Chem. 1976, 68, 313–315, for examples in aromatic oxidations; or (d) Szwarc, M. Acc. Chem. Res. 1972, 5, 169–176 for reductions. 12. See, e.g., Simonet, J.; Lund, H. Acta Chem. Scand. 1977, 31B, 909–911 for a typical experimental example. 13. For a discussion, see Pletcher, D. J. Electroanal. Chem. 1984, 179, 263–267. 14. See, e.g., (a) Lund, H.; Simonet, J. J. Electroanal. Chem. 1975, 65, 205–218 for an early synthetic example. 15. Amatore, C.; Savéant, J.-M. J. Electroanal. Chem. 1981, 123, 203–217. 16. Amatore, C.; M’Halla, F.; Savéant, J.-M. J. Electroanal. Chem. 1981, 123, 219–229. *

Note that this presentation considers only one branching point, but when several are present the same analysis can be applied at each branching point.

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392 17. 18. 19. 20. 21. 22. 23. 24. 25.

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Amatore, C.; Pinson, J.; Savéant, J.-M.; Thiébault, A. J. Electroanal. Chem. 1981, 123, 231–242. Amatore, C.; Savéant, J.-M. J. Electroanal. Chem. 1981, 125, 1–21. Amatore, C.; Savéant, J.-M. J. Electroanal. Chem. 1981, 125, 23–29. Amatore, C.; Savéant, J.-M. J. Electroanal. Chem. 1981, 126, 1–19. Amatore, C.; Savéant, J.-M. J. Am. Chem. Soc. 1981, 103, 5021–5023. See, e.g., Andrieux, C. P.; Savéant, J.-M.; Zann, D. Nouv. J. Chim. 1984, 8, 107–116. Farnia, G.; Severin, M. G.; Capobianco, G.; Vianello, E. J. Chem. Soc. Perkin Trans. II 1978, 1–8. Grimshaw, J.; Hamilton, R.; Trocha-Grimshaw, J. J. Chem. Soc. Perkin Trans. I 1982, 229–234. (a) Amatore, C.; Combellas, C.; Pinson, J.; Savéant, J.-M.; Thiébault, A. J. Chem. Soc. Chem. Commun. 1988, 7–8. (b) Alam, N.; Amatore, C.; Combellas, C.; Pinson, J.; Savéant, J.-M.; Thiébault, A.; Verpeaux, J. N. J. Org. Chem. 1988, 53, 1496–1506. (c) Alam, N.; Amatore, C.; Combellas, C.; Thiebault, A.; Verpeaux, J. N. Tetrahedron Lett. 1987, 28, 6171–6174. 26. M’Halla, F.; Pinson, J.; Savéant, J.-M. J. Am. Chem. Soc. 1980, 102, 4120–4127.

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Section III Electron Transfers and Concerted Processes

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11

Influence of Molecular and Medium Effects on Two-Electron Processes Kevin Lam and William E. Geiger

CONTENTS I. II.

Introduction and Scope ......................................................................................................... 396 Thermodynamic Aspects ...................................................................................................... 396 A. Expression of Potentials ................................................................................................ 396 B. Normal versus Inverted Ordering of E1/2 Potentials ..................................................... 397 C. Free Energy Factors ...................................................................................................... 398 D. Intrinsic versus Extrinsic Factors Affecting ΔE1/2 Values ............................................ 399 E. CV Images for Different ΔE1/2 Values ..........................................................................400 F. Examples of Calculated EE Processes..........................................................................400 1. EE with Minimal Structure Change: Anthracene .................................................400 2. EE with Major Structure Change...........................................................................402 3. Application to Reduction of Cyclooctatetraene .....................................................404 III. Effects of Slow Heterogeneous Electron Transfer ................................................................406 IV. Experimental Aspects: Determination of E.T. Stoichiometry (One or Two Electrons?) ......408 A. Analysis by a Single Electroanalytical Method ............................................................408 B. Analysis by Combined Methods ...................................................................................409 1. Electrolysis and Voltammetry................................................................................409 2. Dual Voltammetric Analysis ................................................................................. 410 V. Applications to Two-Electron Processes .............................................................................. 410 A. Involving a Single Redox Site ....................................................................................... 411 1. Dominant Electronic and Structural Changes ....................................................... 411 2. Combined Structural and Medium Effects ............................................................ 415 3. Dominant Medium Effects .................................................................................... 419 4. Effect of Slow E.T. Kinetics: Reduction of Bis(hexamethylbenzene)ruthenium2+ ....421 B. Involving Two Identical Redox Sites ............................................................................ 422 1. Electronic Communication through the Linkage .................................................. 422 2. Medium Effects...................................................................................................... 425 VI. An Integrated Approach to Medium Effects on ΔE1/2 ......................................................... 427 Acknowledgments.......................................................................................................................... 430 References ...................................................................................................................................... 430

395

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I. INTRODUCTION AND SCOPE Electron transfer (E.T.) reactions in which there is an overall stoichiometry of two or more electrons are common and may be found in a number of important chemical and biological processes. Many of these reactions occur through mechanisms in which chemical reactions (C) are coupled to the E.T. process (E). For example, the classic quinone/hydroquinone couple follows an electrochemical reaction–chemical reaction–electrochemical reaction–chemical reaction mechanism, combining a pair of electron and proton transfers, displaying two-electron Nernstian behavior. This chapter restricts itself to the simpler case of E.T. reactions of two-electron stoichiometry that are uncomplicated by coupled chemical reactions. This will facilitate discussion of the three principal factors affecting the relative ordering and degree of potential separation of the successive one-electron transfers of the Electron transfer/Electron transfer (EE) mechanism (Equation 11.1): changes in electronic and molecular structure, differences in solvation energies, and differences in ion-pairing energies, between members of the E.T. series. Furthermore, only cases in which structure change and medium effects are thought to + e−

+ e−

EA/A −

E − 2− A /A

2− −  ⇀  ⇀ A↽ A ↽ A o o

(11.1)

be concomitant (i.e., concerted) with E.T. will be considered. Thus, “square schemes” [1], in which reversible structure changes “follow” E.T., are not covered. Note also that, although the principles discussed in this chapter relate also to systems having more than two sequential E.T. reactions, “super” multielectron transfer processes such as those involving C60 will not be systematically covered. As will become apparent, a key experimental goal is to distinguish between obviously separated E.T. processes and those which appear to proceed through a single two-electron reaction in the following equation: +2 e −

2−  ⇀ A↽ A o EA/A2−

(11.2)

Writing the E.T. reaction with a two-electron stoichiometry naturally raises the question of whether or not it is possible to have a truly concerted two-electron transfer, that is, a reaction that does not proceed through the one-electron intermediate A−. Whereas it seems clear that a concerted two-electron transfer is theoretically permitted [2], there appear to be no experimentally verified examples of such a process. In evaluating the literature on this subject, Evans concluded that a concerted two-electron process requires that the value of EAo − /A2− be at least 0.4 V (and perhaps as much as 1V) “positive” o − of EA/A − , making disproportionation of the putative one-electron intermediate A energetically so favorable as to limit the experimental ability to differentiate between concerted and separate twoelectron processes [3]. We will treat all E.T. reactions as occurring one electron at a time. Although thermodynamic and kinetic factors are both of obvious importance to the voltammetric behavior of EE reactions, we will give primary attention to the ordinarily more dominant thermodynamic effects. After going over some of the basic thermodynamic aspects of EE reactions and reviewing the electrochemical methods best suited to characterize them, applications are taken from the literature of organic and organometallic redox chemistry.

II.

THERMODyNAMIC ASPECTS

A.

EXPRESSION OF POTENTIALS

The relationships between the three most commonly used terms for the potentials of redox couples, namely, the standard potential (Eo), the formal potential (Eo′), and the E1/2 potential, were discussed in Chapter 1. Referring to Equation 11.3, E1/2 is identical to the standard potential when the activity

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coefficients ( f ) and diffusion coefficients (D) are equal for the two members of the redox couple. We will employ both Eo and E1/2 in the theoretical section but refer exclusively to E1/2 values when experimental examples are discussed. Within these conventions, we employ symbols based on the relative ordering of the redox states of the molecule, with the starting point designated by (0). Thus, reduction of a compound by two one-electron processes (Equation 11.4) takes the fully oxidized o state Ox(0) successively to the first reduced state Red(1) (at potential ERed1 ) and then to the second o reduced state Red(2) (at potential ERed2). A similar nomenclature is employed for successive oxidations of neutral compounds, this time beginning with the neutral compound being in the least oxidized (most reduced) redox state (see Equation 11.5):

E1/ 2

1/ 2 RT  fOx   DRed    =E + ln   nF  fRed   DOx     o

−

−

+e +e  ⇀  ⇀ For reductions: Ox(0) ↽  Red(1) ↽  Red(2) o o ERed1

ERed2

− e−

− e−

EOx1

EOx 2

 ⇀  ⇀ For oxidations: Red(0) ↽  Ox(1) ↽ o  Ox(2) o

(11.3)

(11.4)

(11.5)

An important factor in this chemistry is the magnitude of the potential separations for EE processes, ΔE1/2. So that increasingly positive values of ΔE1/2 always denote systems for which disproportionation of the one-electron product (usually a radical) is increasingly disfavored, the definition of ΔE1/2 is changed for sequences of reductions (Equation 11.6) versus oxidations (Equation 11.7). Thus, a “normal” ordering of the ET processes (i.e., second reduction [“oxidation”] more negative [“positive”] than the first) always has a “positive” ΔE1/2 value and an “inverted” ordering [4] (first reduction [“oxidation”] less negative [“positive”] than the second) always has a “negative” ΔE1/2 value:

B.

1 Red 2 For Reductions: ∆E1/ 2 = E1Red / 2 − E1/ 2

(11.6)

2 Ox 1 For Oxidations: ∆E1/ 2 = E1Ox / 2 − E1/ 2

(11.7)

NORMAL VERSUS INVERTED ORDERING OF E1/2 POTENTIALS

Consider the case of a neutral compound such as anthracene, 1, being reduced to a dianion in a sequence of two E.T. steps. Owing strictly to electrostatic considerations, the second reduction is − /2 − expected to be energetically more difficult than the first, accounting for the fact that E1/2 is lower 0/− (more negative) than E1/2 . This so-called normal ordering is observed in the great majority of overall two-electron processes. The sign of ΔE1/2 has a profound effect on the chemistry of the E.T. series, since it determines whether or not disproportionation of the radical (1− in the present case) (Equation 11.8) is thermodynamically favored. Using Equation 11.9 to obtain Kdisp from ΔE1/2, percentages of the “actual” concentration of the radical in 2 1− ⇌ 1 + 12−

K disp =

 RT ∆E1/ 2 =   F

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[1][12− ] [1− ]2

  ln K disp 

(11.8)

(11.9)

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TAbLE 11.1 Kdisp Calculated for Several ΔE1/2 Values (RT/F = 0.0257 V at 298 K; ΔE1/2 in mV) ΔE1/2

Kdisp

+500 +236 +118 +59 0 −59 −118 −236 −500

3.5 × 10 1.0 × 10−4 1.0 × 10−2 0.10 1 9.9 99 9.7 × 103 2.8 × 108

% Radical Remaining >99.9 98 83 61 33 14 5 0.50 E1/ 2 ), there is only one coupled cathodic/anodic feature (Figure 11.1, ΔE1/2 = 0, −59, and −118 mV). With a ΔE1/2 value lower (more negative) than −118 mV, disproportionation of Red(1) dominates and the reduction of Ox has the experimental appearance of a single two-electron process in the following equation: +2 e −

o  ⇀ Ox(0) ↽  Red(2) E =

1 o o ERed1 + ERed2 2

(

)

(11.13)

Square wave voltammetry scans for the same ΔE1/2 values are shown on the right side of Figure 11.1. Application of this technique will be covered in Section IV.A.

F.

EXAMPLES OF CALCULATED EE PROCESSES

1. EE with Minimal Structure Change: Anthracene In the case of anthracene, 1, the additional gas-phase energy for electron attachment of a second electron is a prodigious 4.2 eV for the electrostatic reasons referred to previously. After accounting for solvation energies calculated by a simple continuum model for a solvent with a dielectric constant

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Influence of Molecular and Medium Effects on Two-Electron Processes

ΔE1/2 (mV)

Cyclic Voltammetry

Square Wave Voltammetry

I (A) 0.0004

I (A)

0.0002

0.0004

0

0

–0.0002

–0.0004

236

–0.0004

–0.0008 –0.6

–0.8

–1

–1.2

–1.4

–1.6

–1.8

–2

–1

–1.1

–1.2

–1.3

E (V)

–1.8

–1.5

–1.6

–1.7 –1.8

–1.5

–1.6

–1.7 –1.8

–1.5

–1.6

–1.7 –1.8

–1.5

–1.6

–1.7 –1.8

–0.0004

–0.0004 –0.6

–0.8

–1

–1.2 –1.4 E (V)

–1.6

–1.8

–2

–0.0008 –1

–1.1

–1.2

–1.3

–1.4 E (V)

I (A) 0.001

I (A) 0.0006 0.0004 0.0002 0 –0.0002 –0.0004 –0.0006

0.0005 0 –0.0005 –0.001 –0.6

–0.8

–1

–1.2

–1.4

–1.6

–1.8

–1

–2

–1.1

–1.2

–1.3

E (V)

–1.4 E (V)

I (A)

I (A) 0.0008

0.001

0.0004

0

0

–0.001

–0.0004 –0.0008 –0.6

–0.8

–1

–1.2

–1.4

–1.6

–1.8

–2

–0.002 –1

–1.1

–1.2

–1.3

–1.4 E (V)

E (V) I (A) 0.0008

I (A) 0.002

0.0004

0.001 0

0

–0.001

–0.0004

–0.002

–0.0008 –0.6

–0.8

–1

–1.2

–1.4

–1.6

–1.8

–2

–1

–1.1

–1.2

–1.3

E (V)

–1.4 E (V)

I (A) 0.002

I (A) 0.0008

0.001

0.0004

–118

–1.7

0

0

118 –0.0002

–59

–1.6

0.0004

0.0002

0

–1.5

I (A)

I (A) 0.0004

59

–1.4 E (V)

0

0

–0.001

–0.0004

–0.002

–0.0008 –0.6

–0.8

–1

–1.2 E (V)

–1.4

–1.6

–1.8

–2

–1

–1.1

–1.2

–1.3

–1.4

–1.5

–1.6

–1.7 –1.8

E (V)

FIgURE 11.1 Simulated (DigiElch) cyclic and SWVs for two consecutive reversible one-electron reductions for different ΔE1/2. Parameters for cyclic voltammograms: v = 200 mV s−1, k0 = 104 cm s−1, α = 0.5, T = 298.2 K, CA* = 1 mM, and A = 1 cm². For square wave, f = 100 Hz.

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of 36 (mimicking CH3CN or DMF), the calculated potential separation was reduced to 0.87 eV [8] (calculationally, we have now reached the middle of Scheme 11.1). Inclusion of estimated ion-pairing energies (not included in Scheme 11.1) of [NMe4]+ ions with 1− and 12− further lower the calculated potential separation to 0.83 V, reasonably close to the measured ΔE1/2 value of 0.67 V [9]. The fact that the ΔE1/2 values decrease as the medium effects are taken into account is understood intuitively by increases in both Born solvation energies and coulombic ion-pairing strengths with increasing charge of the products. However, the sum of the intrinsic and extrinsic stabilization energies gained Red1 in formation of the anthracene dianion is insufficient to bring E1Red2 / 2 positive of E1/ 2 . Referring back to Scheme 11.1, the intrinsic structural relaxations specified in the bottom section are small. The reduction of anthracene therefore obeys normal potential ordering, with a positive value of ΔE1/2. A chemical consequence of the ΔE1/2 value for the 1/1−/12− ET series is that, given the absence of a follow-up reaction such as protonation, one-electron reduction of anthracene, whether accomplished strictly electrochemically or with a chemical reducing agent, gives a stable solution of its anion radical. Specifically, the experimental ΔE1/2 value of 0.67 V [9] gives a Kdisp value of 5 × 10 −12, and much smaller values of ΔE1/2 are necessary for there to be analytically or chemically significant amounts of the disproportionation products 1 and 12− (see Table 11.1). 2. EE with Major Structure Change In the great majority of systems exhibiting inverted EE behavior, a major structural change accompanies at least one of the E.T. reactions. It is important to reiterate the distinction between this case and “square schemes.” The latter are, in fact, sequences of electrochemical–chemical reactions [1,3,10]. As discussed in Section 1, we will consider only electrochemical reactions, for which structure change is considered to be concomitant with E.T. a. Structure Change in Second ET Process Within the EE model (Equations 11.14 and 11.15), consider the case in which the structure change occurs in the second E.T. process. In Equation 11.15, the symbol “*” indicates that the molecule has undergone a major structural rearrangement. Referring again to the bottom row of Scheme 11.1, the minor versus major structural relaxation energies + e−

(11.14)

Ox/Red1 ⇀ Ox(0) ↽  Red(1) E1/ 2

+ e−

(11.15)

* Red1/Red*2 ⇀ Red(1) ↽  Red (2) E1/ 2

of the two E.T. steps cause a more positive shift in the second E.T. process compared to the first. If the structural stabilization energy is sufficient, ΔE1/2 may be negative (i.e., inverted), and the reduction of Ox(0) follows the two-electron stoichiometry of Equation 11.16. We will treat examples of such a system in the following text when considering the reductions of bis(hexamethylbenzene) ruthenium (II) and arene cyclopentadienyl iridium (III) complexes: −

+2 e *  ⇀ Ox(0) ↽  Red (2) E1/ 2 =

1 Ox/Red1 2 E1/ 2 + E1Red1/Red* /2 2

(

)

(11.16)

b. Structure Change in First E.T. Process That structural change in the second E.T. step may lead to potential inversion is intellectually straightforward, in the sense that the intrinsic and extrinsic stabilization forces (structure change and medium effects, respectively) are conveying the same directional sign changes on ΔE1/2, each imparting a positive shift of E1/2(2) relative to E1/2(1). Less intuitive is the fact that structure change

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Influence of Molecular and Medium Effects on Two-Electron Processes

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Ox/Red*1 > Red*1/Red*2 E

Large relaxation energy for monoanion, ΔE1/2 positive

B

(normal ordering)

A

B

C

Ox*/Red*1 Red*1/Red*2

Ox/Red1

E

B

D

Red1/Red2

Small relaxation energy for monoanion, ΔE1/2 negative (inverted ordering)

Red*1/Red*2 > Ox/Red*1 More positive E1/2

SCHEME 11.2 Effects of structural relaxation in first reduction on ordering of potentials for EE reductions. The middle section shows the medium-adjusted EE potentials for the two different “frozen” structures designated as either unmarked or with a “*.” The labels A through E are discussed in the text. Solid arrow used for E.T. involving major structure change.

in the first E.T. step may also give rise to inverted ΔE1/2 potentials. For a reduction, this requires that the rearranged (*) structure have dramatically better electron acceptor properties than the original structure, so that the second reduction of putative Ox*(0) would be more facile than the first reduction of Ox(0) to Red*(1). Consider Scheme 11.2, which begins with the assumption that account has already been taken of the extrinsic (electrolyte) effects. The middle of the scheme pictures a pair of hypothetical EE processes for two structurally different molecules assumed to be “locked” into their two different isomeric structures. Their two individual E.T. series (Equations 11.17 and 11.18) are represented by the less + e−

+ e−

 ⇀  ⇀ Ox(0) ↽  Red(1) ↽  Red(2) Ox/Red1 Red1/Red2 E1/ 2

(C )

E1/ 2

+ e−

(D )

+ e−

 ⇀  ⇀ Ox * (0) ↽  Red*(1) ↽  Red*(2) Ox*/Red*1 Red*1/Red*2 E1/ 2

(A )

E1/ 2

(B )

(11.17)

(11.18)

negative pair of lines, A and B, for Ox*(0) and the more negative pair, C and D, for Ox(0), the latter being the true starting material. The upper and lower drawings show the possible consequences of introducing a major structure change into the first reduction of Ox(0). In both cases, the couple Ox(0)/Red(1) (C) becomes Ox(0)/Red*(1) (E), modifying the potential of the first reduction by the amounts of the solid arrows. The potential E1Red*1/Red*2 (B) remains unaffected and those of /2 (D) are no longer pertinent. (A) and E1Red1/Red2 E1Ox*/Red*1 /2 /2

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If the stabilization energy shifts E positive of B (top drawing), the system retains normal potential ordering, with its positive ΔE1/2 value now representing E1Ox/Red*1 . If the structural − E1Red*1/Red*2 /2 /2 stabilization energy is small (bottom drawing), then B resides positive of E, inverting the E1/2 values and leading to a single two-electron wave as Ox(0) goes to Red*(2) (Equation 11.16). One example of a system that can be experimentally manipulated to exhibit either normal or inverted behavior is the reduction of cyclooctatetraene, which we now consider. 3. Application to Reduction of Cyclooctatetraene Reduction of tub-like COT to its planar dianion (COT2−) has been intensely studied through a number of chemical, electrochemical, and theoretical methods [11], with general agreement that a two-step EE mechanism is involved. Studies under dissimilar experimental conditions have shown that this system switches between normal and inverted EE behavior, depending on the electrolyte medium. In most electrolytes, COT undergoes two normally ordered reductions having potentials that are closely spaced, at least relative to anthracene and other aromatic hydrocarbons. Putting aside for the moment the effects of slow heterogeneous E.T. on the shape of the first waves, the d.c. and a.c. polarographic scans of Figure 11.2 clearly demonstrate the normally ordered EE process, COT/COT−/COT2−, in DMF/0.1 M [NBu4][ClO4] (ΔE1/2 = +0.24 V) [12]. However, an inverted ΔE1/2 value of −0.22 V was estimated for the same redox process in liquid ammonia (0.1 M KI) [13], in which the CV scan (Figure 11.3) shows only a single matched pair of cathodic and anodic features consistent with a two-electron process, COT/COT2−. +e–

+e– E1/2

0/–

–

2–

E1/2–/2–

COT2–

COT–

COT

A quantitative treatment of the thermodynamics of this system was detailed by Baik et al. [14]. Scheme 11.3 demonstrates the principal energetic impacts of structural relaxation and ion pairing on the potentials of the EE sequence for COT. Conceptually, there are four redox couples that must be accounted for, based on two idealized structures, either tub (T0/−/2−) or planar (P0/−/2−). (We are simplifying this, considering only a delocalized dianion. In fact, the calculations also take into account the energy gained by the planar anions going from non-delocalized to delocalized electronic structures.) The computed potentials sketched in the top row of Scheme 11.3 show that reductions of both isomers would follow normal EE mechanisms if locked into their given tub or 1.50

4.00

1.00 μA

μA

3.00 2.00

0.50 1.00 0.00 –1.00 (a)

–1.10

–1.20

–1.30

–1.40

Applied d.c. potential (mV)

0.00 –1.00

–1.50 (b)

–1.10

–1.20

–1.30

–1.40

Applied d.c. potential (mV)

FIgURE 11.2 Comparison of theory and experiment for the first two polarographic waves of COT. (o) Experimental d.c. current (a) and experimental fundamental harmonic inphase a.c. polarographic current (b). (⎯) theoretical polarograms for ks1 = 2.0 × 10 − 3 cm s−1, α1 = 0.40, ks2 = 0.15 cm s−1, α2 = 0.50, and Do = 1.4 × 10 −5 cm2 s−1. (From Huebert, B.J. and Smith, D.E., J. Electroanal. Chem., 31, 333, 1971. With permission.)

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Influence of Molecular and Medium Effects on Two-Electron Processes

–0.8

–1.2 E(V) vs. Ag/Ag+

405

–1.6

FIgURE 11.3 Simulated voltammogram of COT (- - -) versus experimental voltammogram (–). ks1 = 1.58 × 0 0 10 −4 cm s−1, ks2 = 1.26 × 10 −2 cm s−1, (E1 − E2 ) = − 220 mV, T = 245 K. (From Smith, W.H. and Bard, A.J., J. Electroanal. Chem., 76, 19, 1977. With permission.)

planar forms, with ΔE1/2 values of about +0.76 V for the planar isomer and 1.05 V for the tub isomer. Importantly, both E1/2 potentials of the planar form are higher (more positive) than the first potential of the tub form. Introduction of structural relaxation into the calculations (middle sketch) shifts the reduction of the neutral tub positively by 1.02 eV as the monoanion takes on the planar geometry. E1/2(1) now refers to the couple T0/P−. The second reduction, E1/2(2), is necessarily that of P−/P2−, which shifts only slightly owing to the much smaller relaxation energy (0.13 V) inherent to formation of the planar dianion from the structurally similar monoanion. Referring back to our earlier theoretical treatment, the COT system is now described by the top sketch of Scheme 11.2. The calculated ΔE1/2 value of + 0.41 V does not take into account the modest ion-pairing effects of [NBu4]+ with COT2−, which would likely reduce this value by about 0.1 V and bring it very close to the measured value [12] of +0.24 V. The strong ion pairing required in order to bring E1/2(2) positive of E1/2(1) was quantitatively evaluated by introducing a solvated ion-pair model based on potassium into the calculations (bottom sketch of Scheme 11.3). The strong ion pairing of K2(COT) provides the final stabilizing energy to achieve potential inversion, with the calculated ΔE1/2 of −0.25 V being remarkably close to the experimental value [13] of −0.22 V in NH3(l)/0.1 M KI.

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T0/T–

P–/P2–

0.76 V

0.13 V

T0/P–

P–/P2–

0.73 V

1.05 V

Solvated, rigid structures P = planar, T = tub

1.02 V

P–/P2–

T0/P–

T–T2–

Introduce structural relaxation, ΔE1/2 = +0.41 V (normal)

Introduce K+ ion-pairing ΔE1/2 = –0.25 V (inverted)

More positive E1/2

SCHEME 11.3 Sequence of calculated energies of E1/2 values of COT reductions, beginning from solvated species having either a planar (P) or tub (T) geometry.

III.

EFFECTS OF SLOW HETEROgENEOUS ELECTRON TRANSFER

It is well known that a structure change such as the COT tub-to-planar rearrangement that occurs concomitant with E.T. lowers the standard heterogeneous E.T. rate ko (sometimes labeled ks) owing to an increase in the inner-shell reorganization energy λi of the reaction. The redox couple then behaves as a quasi-reversible or electrochemically irreversible E.T. system. Given that structural reorganization is often a key factor in inducing inverted potential ordering, it is not surprising that sluggish E.T. kinetics are often observed with inverted two-electron systems. Furthermore, the effects of slow E.T. on voltammetric displays are more notable than in the normally ordered case. Figure 11.4 contains CV scans for a normally ordered EE system (ΔE1/2 = +200 mV) in which (A) both E.T. steps are Nernstian (fast); (B) the first E.T. is quasi-reversible (slow); and (C) the second E.T. is quasi-reversible. The “spread out” appearance of the first wave in case B has its counterparts in other voltammetric methods, as seen in the dc and ac polarographic scans of Figure 11.2. Figure 11.5 gives the same sequence of fast and slow E.T. for an inverted EE system with ΔE1/2 = −200 mV. It is immediately obvious that, compared to the normally ordered case, sluggish E.T. for an inverted system has a much more dramatic effect on the shapes and positions of the CV waves. In general, the forward peak (cathodic in this case) appears close to the potential of the first E.T. step, whereas the reverse (anodic) peak appears close to the potential of the second E.T. step. When both E.T. steps are fast (D), the cathodic and anodic peaks appear near the average of the two E1/2 potentials. The qualitative reason for this effect has to do with how one slow (i.e., rate determining) charge transfer affects the concentration of the other E.T. product at given applied potentials. Consider, for example, scans D (both E.T. fast) and E (first E.T. slow) in Figure 11.5. With the measured E1/2 of scan D being the average of E1/2(1) and E1/2(2), (0.1 V here), the first reduction of Ox(0) begins at a significant underpotential, that is, quite positive, of E1/2(1). However, Red(1) is produced at a significant overpotential for its reduction to Red(2). Thus, as long as both E.T. steps have high ko values, the system moves quickly to equilibrium, and the current/potential curve is governed by the Nernst equation and mass transfer. The resulting forward wave is Nernstian shaped. However, if the E.T. of the first reduction is slow (scan E), it becomes rate determining and governs the current/potential plot.

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Influence of Molecular and Medium Effects on Two-Electron Processes 2.00E–03 1.50E–03 A

Current (mA)

1.00E–03

C

5.00E–04

B

0.00E+00 –5.00E–04 ks1= 1 ks2 = 1 ks1 = 0.02 ks2 = 1 ks1 = 1 ks2 = 0.01

–1.00E–03 –1.50E–03 0.3

0.2

0.1

0

–0.1 –0.2 Potential (V)

–0.3

–0.4

–0.5

–0.6

FIgURE 11.4 Simulated (DigiSim) cyclic voltammograms for normal ordered two consecutive reversible one-electron reductions with different heterogeneous standard heterogeneous E.T. rate constants in cm s−1. Parameters: E1 = 0 V, E2 = −0.2 V, v = 1 V s−1, α = 0.5, T = 298.2 K, CA* = 1 mM, and A = 1 cm². 4.00E–03 3.00E–03

F

D

E

Current (mA)

2.00E–03 1.00E–03 0.00E+00

–1.00E–03

ks1 = 1 ks2 = 1 ks1 = 0.02 ks2 = 1 ks1 = 1 ks2 = 0.01

–2.00E–03 –3.00E–03 0.6

0.5

0.4

0.3

0.2 0.1 Potential (V)

0

–0.1

–0.2

–0.3

FIgURE 11.5 Simulated (DigiSim) cyclic voltammograms for inverted two consecutive reversible oneelectron reductions with different heterogeneous standard heterogeneous E.T. rate constants in cm s−1. Parameters: E1 = 0 V, E2 = +0.2 V, v = 1 V s−1, α = 0.5, T = 298.2 K, CA* = 1 mM, and A = 1 cm².

Thus, when the potential reaches 0 V, a diminished current is obtained owing simply to the slow conversion of Ox(0) to Red(1). The expected currents are not observed until the applied potential becomes close to E1/2(1), whereupon the forward wave takes on the shape of an irreversible cathodic process. In the alternate case in which the second reduction is rate determining (scan F), the cathodic wave again appears near E1/2(1), but it is more “reversibly” shaped because, when significant quantities of Red(1) are produced, the applied potential is already at a high overpotential for the reduction of Red(1) to Red(2). These arguments have been well expressed by Evans [15].

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Whereas the CVs in Figures 11.4 and 11.5 were produced using transmission coefficients, α, of 0.5, it should be mentioned that the shape of a quasi-reversible wave of either one-electron or twoelectron type is quite sensitive to the α value. The fact that E.T. reactions in which α ≠ 0.5 are fairly common in two-electron chemistry means that the normal CV diagnostics (Ep − Ep/2; ip/v1/2, Ep shifts with v, etc.) are often not easily applied to such systems. For inverted potential systems, computer simulations of CV curves at multiple scan rates are almost always necessary to extract anything more than qualitative conclusions about the individual one-electron steps.

IV. EXPERIMENTAL ASPECTS: DETERMINATION OF E.T. STOICHIOMETRy (ONE OR TWO ELECTRONS?) When ΔE1/2 values are less than about 60 mV, the fact that the voltammetry shows only a single cathodic/anodic pair (see Figure 11.1) introduces the need for additional analysis to tell whether the redox process has a stoichiometry of one or two electrons. Several different approaches to this problem are given in this section. Note that the voltammetric approaches refer only to two-electron processes involving a single-site redox system and exclude the kinds of multisite systems that might be found in polymers, dendrimers, and molecules that are “tagged” with equivalent redox-active groups such as ferrocenyl moieties (Section V). Also, the discussion is restricted to simple EE systems without coupled reactions.

A. ANALYSIS BY A SINGLE ELECTROANALYTICAL METHOD For a diffusion-controlled system, the voltammetric current is a function of both the n-value of the couple and the diffusion coefficient, Do, of the analyte. The latter is often estimated by using the known Do value of a molecule of similar shape and size in the same medium. The diffusion coefficient for an uncharged spherical molecule is given by the Stokes–Einstein equation (Do = RT/6NAπηr), where NA is Avogadro’s number, η is the solvent viscosity, and r is the spherical radius. For spheres of different molecular weights, M, Do ∝ M0.33 and for linear molecules, Do ∝ M0.55 [16]. Recently, an NMR method has been introduced as a convenient, electrochemically independent approach to the determination of Do values [17,18]. However, even if Do is known, or can be acceptably estimated, complications arise in the use of CV peak currents to determine n values. Consider the following: The current function, χ, is defined by Equation 11.19, where ip is the peak current in μA, v is the scan rate in V s−1, n is the number of electrons transferred, A is the area of the electrode in cm2, Do is the diffusion coefficient of the electroactive species in cm2 s−1, and Co is its bulk concentration in mol cm−3 [19]. Its prediction that the scan-rate normalized peak current for a two-electron process is 2.83 times that of a one-electron process holds only when the second oneelectron process has a much milder E1/2 x=

ip = 0.269n3 / 2 ADo1/ 2C ° v1/ 2

(11.19)

potential than the first one-electron process. Considering reductions, for example, it requires that be significantly more positive than E1Ox/Red1 . If the two E1/2 values are closer, the cathodic E1Red1/Red2 /2 /2 peak height becomes a sensitive function of ΔE1/2 and digital simulations are needed to solve the question of the E.T. stoichiometry. The shape of the CV wave is also sensitive to the E.T. stoichiometry. The most readily measured CV parameter is the separation between cathodic and the anodic current peaks, Epc and Epa, respectively, which follows the relationship in the following equation: Epa − Epc ≈ RT l n(10)/nF (≈ 60 mV/n at room temperature)

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(11.20)

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Influence of Molecular and Medium Effects on Two-Electron Processes

TAbLE 11.2 Partial Listing of the Effects of Changes in ΔE1/2 on the CV ΔEp (= Epa − Epc) Value and the DPV Width at Half Height (W1/2) Mechanism E (single one electron) EE EE EE EE EE

ΔE1/2 (mV)

ΔEp (mV)

n.a. +118 (normal) +59 (normal) 0 −59 (inverted) −118 (inverted)

60 ± 1 161 84 42 34 30

a

W1/2 (mV) (DPV) 90 207 122 66 52 48

Nicholson, R.S., Anal. Chem., 37, 1351, 1965; Nicholson, R.S. and Shain, I., Anal. Chem., 36, 706, 1964. Values for EE system are taken from Reference 22. The calculations assume Nernstian E.T. behavior and the DPV data are based on a 10 mV pulse amplitude. DPV stands for differential pulse voltammetry. a The ΔE value for an E mechanism is sensitive to the difference between E and the experimental p pc switching potential, Eλ.

Sources:

Thus, a ΔEp value of significantly less than 60 mV is a clear indication of an inverted two-electron system [16,20,21]. Conversely, if normally ordered ΔE1/2 values are small enough to preclude resolution of sequential waves, the single forward and reverse waves are broader than those of a Nernstian one-electron process (see ΔE1/2 = +59 mV in Figure 11.1). Digital simulation of an experimental CV wave could be employed to obtain the ΔE1/2 value, but published tables can serve this purpose in most cases. Particularly useful are data published by Richardson and Taube [22], examples of which are collected in Table 11.2. Note that the CV peak separation is 42 mV when E1/2(2) = E1/2(1) and only achieves the “true” two-electron value predicted by Equation 11.20 when the degree of potential inversion reaches 118 mV. The ΔE1/2 values for overlapped EE waves can also be determined by measuring the widths at half height, W1/2, obtained using pulsed voltammetric methods. Simulations of square wave voltammograms (SWVs) for EE systems are shown in Figure 11.1. Similarly shaped waves are obtained by differential pulse voltammetry (DPV) and have the advantage of being less sensitive to sluggish charge-transfer kinetics. Equation 11.21 indicates that when the DPV pulse height (ΔEpulse) is kept small, the half width is about 90 mV for a one-electron process and 45 mV for a two-electron process. Calculated W1/2 values for representative ΔE1/2 values are given in Table 11.2, and a more exhaustive working curve is available in Reference 22: for ∆Epulse < 20 /n mV, W1/ 2 = 90 /n mV

(11.21)

B. ANALYSIS BY COMBINED METHODS 1. Electrolysis and Voltammetry Controlled potential electrolysis, in which one measures the coulometry required for a complete bulk redox reaction, is a powerful method for determining the n-value of an electrochemical process. Within the present context, a few cautionary notes are in order. Primary among these is the difference in timescale between voltammetry (usually less than 10 s) and bulk electrolysis. With the exception of experiments carried out in thin-layer cells, bulk electrolyses typically take minutes to scores-of-minutes to complete, raising the possibility that the short-time and long-time E.T. mechanisms and products might be different. Voltammetric analysis before and after exhaustive

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electrolysis may minimize the chances of undetected timescale-dependent changes in mechanism, but the micro- versus macroscale natures of the different techniques must always be kept in mind. One should also be aware that, owing to ohmic drop, the effective potential may vary considerably from the nominal Eappl across the surface of a large electrode when the bulk electrolysis is carried out in a nonaqueous electrolyte solution. This can introduce a considerable problem when closely spaced reactions are involved. A review of how coupled chemical reactions affect bulk coulometric experiments is available [23]. 2. Dual Voltammetric Analysis Based on the preceding discussion, there is a clear need for a way to determine the E.T. stoichiometry that is not sensitive to differences in experimental timescales and is independent of the diffusion coefficient of the test compound. This can be accomplished by using the results of two electroanalytical experiments, one based on a “transient” technique (e.g., CV or chronoamperometry) and the other on a “steady-state” technique (e.g., rotating disk voltammetry or ultramicroelectrode voltammetry). The fundamental idea, apparently first suggested by Lingane [24], involves measuring technique-sensitive characteristic currents that differ in their proportionality to the diffusion coefficient of the test compound. For example, consider that the fundamental measurable in chronoamperometry (the it1/2 value) is proportional to nDo1/2 (Cottrell equation, it 1/ 2 = n /π1/ 2 FADo1/ 2Co ), but those of the steady-state methods are proportional to either nDo2/3 (through the Levich equation, ilim /ω1/ 2Co = 0.62 nFADo2 / 3ν −1/ 6, where ω = angular rotation rate and ν = kinematic viscosity of solution) or nDo (for hemispherical diffusion to ultramicroelectrode disk of radius r, ilim = 4nFDorCo). Taking as an example the reduction of COT in liquid ammonia (Figure 11.3), one might obtain chronoamperometric it1/2 data by pulsing from −0.7 to −1.7 V at a disk electrode of 2 mm diameter and then obtain limiting current from a steady-state voltammogram over the same potential range with a 10 μm ultramicroelectrode. The ratio of the two measurables allows the determination of Do, and thereby n. Chronoamperometry has advantages over CV as the transient method, as discussed elsewhere [25].

(

)

V. APPLICATIONS TO TWO-ELECTRON PROCESSES Here, we consider EE applications as involving two different categories of molecules: those having a single redox site and those having multiple, chemically identical, redox sites. Although some structures blur this formal separation, the grouping is pedagogically useful. Certainly, single-site and multiple-site systems may have moieties that are capable of being E.T. sites in and of themselves. What will distinguish the two cases in our treatment is whether or not the “multiple sites” are significant parts of a delocalized molecular redox orbital (i.e., highest occupied molecular orbital [HOMO] for oxidations, lowest unoccupied molecular orbital [LUMO] for reductions). Thus, the intrinsically delocalized molecule bis(fulvalene)dinickel, 2, is treated as a single-site system, but bis(ferrocenyl)ethane, 3, is treated as a two-site system owing to only weak electronic communication between the two ferrocenyl redox centers.

Fe Ni

Ni Fe

2

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Influence of Molecular and Medium Effects on Two-Electron Processes

A.

411

INVOLVING A SINGLE REDOX SITE

1. Dominant Electronic and Structural Changes a. Reduction of [M(η5−C5Me5)(η6−C6Me6)]2+ The net two-electron reduction of the mixed sandwich complexes [MCp*(η6 −C6Me6)]2+, Cp* = (η5−C5Me5), M = Co, Rh, Ir, provides an example of how intrinsic structural and electronic molecular properties can play a major role in determining the ΔE1/2 values of an EE system. The following equation, which defines the E.T. + e−

+ e−

1/ 2

1/ 2

∗ + ∗ m m ⇀ ⇀ [MCp∗ (η6 − C6 Me6 )]2 + ↽  MCp (η − C6 Me6 )  [ MCp (η − C6 Me6 )] ↽ E (1) E (2)

(11.22)

series for these Group 9 metals, denotes the hapticity of the hexamethylbenzene (hmb) ring as η6 for the 18-electron dications (confirmed spectroscopically or crystallographically for all three metals), but does not specify the hapticity of hmb in the +1 and neutral complexes. In fact, as detailed elsewhere [26], the Co complex appears to retain the planar, η6-coordinated, hmb ring throughout the E.T. series, as does the Rh complex in the first reduction to 4+ [27a]. However, arene bending occurs in the second E.T. step of the Rh (and most likely, also Ir) complex, with the structural de-hinging providing an η4-coordinated hmb ring and an electronically more favorable 18 e− structure to neutral 4. The three different metal complexes together offer an informative view of how the intrinsic molecular properties of a congeneric series of metal complexes may affect an EE E.T. mechanism. +

2+

+e–

+e– Rh

42+

Rh

Rh

4+

4

Figure 11.6 gives representative CV scans of the three Group 9 complexes, in which different directions for the shifts of E1/2(1) and E1/2(2) are apparent. In going down from first-row Co to thirdrow Ir, the potential of E1/2(2) shifts markedly positive with respect to E1/2(1), reaching inverted potential behavior for the Ir system, as described by the single two-electron reaction of the following equation [27]: +2 e −

∗ 6  ⇀  [ IrCp∗ (η6 − C6 Me6 )]2 + ↽  IrCp (η − C6 Me6 ) E =[ E (1) + E (2) ]/ 2 1/ 2

1/ 2

1/ 2

(11.23)

The negative shift of E1/2(1) is readily ascribed to an intrinsic electronic effect, as it tracks changes in electron affinities in going from Co to Rh to Ir. The positive shift of E1/2(2), on the other hand, is ascribed to structural relaxation energies that are considerable owing to the planar-to-bent arene rearrangement that accompanies only the second reduction of the Rh and Ir complexes [27]. Another example of the effect of planar-to-bent arene structure change in an EE system is that of the ruthenium sandwich compound [Ru(hmb)2]2+/+/0. Owing to the fact that this system is complicated by slow heterogeneous E.T. kinetics, we will come back to it in Section V.A.4. b. Reductions of Dinitrobenzenes and Dinitrobutenes Although the reductions of most dinitrobenzene derivatives exhibit a normal ordering of potentials, dinitrodurene 5 shows a significant potential inversion (ΔE = −280 mV) [28]. This behavior has been attributed to a major structural change, which makes the radical anion a better electron acceptor

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Organic Electrochemistry E1/2(1)

E1/2(2)

ΔE1/2

M

Co +970 mV Me6 M2+ Me5 Rh + 200 mV

Ir –310 mV

–0.5

0

FIgURE 11.6

–1.0 V vs. SCE

–1.5

Cyclic voltammograms of [M(η5−C5Me5)(η6 −C6Me6)]2+/+/0, M = Co, Rh, Ir.

than the neutral compound 5 with the oxygen atoms of the nitro groups pointing out of the plane owing to steric factors. In the dianion, the nitro groups are sp2 hybridized and have to be in the same plane as the ring. As a consequence, in order to reduce the steric hindrance between the nitro and the methyl groups, the six-membered ring adopts a distorted boat conformation. Evans et al. showed that the standard heterogeneous E.T. rate for the second reduction (0.01 cm s−1) was slower than that of the first (0.20 cm s−1), and accounted for this fact by calculations, which showed that the major structure change (from twisted to folded) occurred in the second reduction [28]. O O

O N

+2e–

N

O

O O N

O N

O 5

52–

Another example of the role of structure change in inducing two-electron behavior in dinitro compounds involves the reduction of trans-2,3-dinitro-2-butene (6) (ΔE 0 = −80 mV) [29], Scheme 11.4. The important structural parameters for the three redox states of 6 are collected in Table 11.3 [1]. Surprisingly, none of the redox states of the dinitrobutene exhibits a fully planar structure. In the neutral form, even if the dihedral angle between C2C1C1C2 is 174.6°, the two nitro groups are significantly tilted out of the plane. In the radical anion, the olefinic bond elongates (Scheme 11.4) and – –

+

O

O

O–N

–

+e

–

O N

–

+e

+

O

O–N

–

N–O O

+

N O O

N–O O

+

–

–

6

6–

62–

SCHEME 11.4 Reduction of trans-2,3-dinitro-2-butene. (From Evans, D.H. and Hu, K.J., Chem. Soc., Faraday Trans., 92, 3983, 1996.)

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Influence of Molecular and Medium Effects on Two-Electron Processes

TAbLE 11.3 Relevant Structural Parameters Calculated by DFT Method for Three Redox States of trans-2,3-Dinitro-2-butene (6) Dihedral Angle (°) Species Neutral (6) Radical anion (6−) Dianion (62−)

bond Length (Å)

C2C1C1C2

O1NC1C1

C1C1

NC1

174.6 150.5 130.0

129.3 168.2 169.9

1.34 1.40 1.46

1.49 1.41 1.34

Source: Evans, D.H. and Busch, R.W., J. Am. Chem. Soc., 104, 5057, 1982.

θ LUMO

H H

H H

HOMO

90°

0° θ

FIgURE 11.7 Qualitative variation of LUMO and HOMO energies of ethylene as a function of the dihedral angle. (From Evans, D.H. and Busch, R.W., J. Am. Chem. Soc., 104, 5057, 1982. With permission.)

twists, while the nitro groups turn into the plane. Finally, in the dianion, the nitro groups are almost coplanar with C2C1C1, but the C1C1 bond twists radically and adopts almost single-bond length. Computations allowed the determination of the difference in the free energy of formation for 6 − and 62−. For the radical-anion 6 −, only a marginal difference in energy is found between a model where the radical anion adopts the same geometry as the neutral compound and a model that uses an optimized geometry: ∆∆Gfo ( g ) = 7.6 kcal mol−1. However, for the dianion 62− the difference is significant: ∆∆Gfo ( g ) = 57.6 kcal mol−1, implying that the major structural reorganization, and origin of the potential inversion, occurs in the second reduction. A qualitative understanding of how the distortion of an olefin can affect its reduction potential is obtained by considering the energy of the frontier orbitals of ethylene as a function of the dihedral angle of the molecule (Figure 11.7). At dihedral angle of 0°, the ethylene moiety is undistorted and the two p orbitals overlap well. Increasing the torsional angle to 90° eliminates the orbital overlap and produces a quasi-degenerate pair. Thus, the energy of the LUMO decreases as the twisting of the molecule increases, accounting for the greater electron affinity of the more highly twisted isomer. c. Oxidation of Tetrathiofulvalenes Another organic example of an inverted potential ordering is found in the oxidation of tetrathiafulvalene (TTF) derivatives, which may be oxidized to their corresponding dications [30]. Depending on the length of the linker between the two dithiafulvalene moieties, the oxidation process may be either a single two-electron process or separate monoelectronic processes. Long chains favor the

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Organic Electrochemistry

(a)

(b)

(c)

FIgURE 11.8 Optimized geometries calculated at the B3LYP/6–31G* level of the neutral (a), radical anion (b), dianion (c) of 8. (From Bellec, N. et al., J. Phys. Chem. A, 104, 9750, 2000. With permission.)

former while short chains favor the latter. Intuitively, the repulsion between the two electrogenerated positive charges is less significant when a long linker is used, a point we will return to in Section V.B. However, in the phenyl-substituted TTF derivative 7, a single 2e − wave is observed in certain media [31]. Thus, whereas acetonitrile favors the 2e− process, a less polar and more donating solvent such as dichloromethane favors two independent 1e− transfers. X-ray crystallography, spectroelectrochemistry, and molecular modeling allow an understanding of this system as one in which dominant structure change accompanies the first redox process (as discussed in Section II.F.2.B). Me

S

MeS

S

S

SMe

S

Me

7

As seen in Figure 11.8, both 7− and 72− adopt a conformation where the TTF cores are almost planar, in contrast to the neutral form 7, in which they are twisted. Calculations delineate the cause of the potential inversion by showing why the radical-anion 7− is a much better electron donor than the neutral compound 7: the energy of the singly occupied molecular orbital of the radical anion is actually raised by the structural changes and as a result, the second reduction becomes easier than the first [31]. d. Reduction of a Carbene-CS2 Adduct As with TTF, a potential compression or even a potential inversion may be observed in the reduction of some N,N′-dialkyl-4,5-dimethylimidazolium-2-dithiocarboxylates (8) (Figure 11.9) [32]. Increasing the bulkiness of the alkyl R substituents on the nitrogens of the carbene moiety, 8, decreases the separation between the potentials of the first and the second reduction (ΔER=Me  =  170  mV, ΔER=Et = 130 mV). In an extreme case, when R = isopropyl, a potential inversion is observed (ΔER=iPr = −30 mV). In this case also, X-ray and NMR data show that a major structural change occurs in the first reduction. Models for the structural changes occurring during the redox processes are depicted in Scheme 11.5. In neutral 8, the dithiolate group is perpendicular to the imidazolium ring. However, during the first reduction, the dithiolate moiety is partially tilted, approaching the imidazolium plane. Finally, reduction to the dianion 82−, which can exist either in triplet state (Scheme 11.5 upper pathway) or a singlet state, results in the formation of an exomethylene moiety (Scheme 11.5 lower pathway). Based on computational studies, the formation of the exocyclic double bond is thermodynamically more favorable than the formation of a biradical. Thus, the formation of the dianion most likely follows a singlet-state pathway [32].

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Influence of Molecular and Medium Effects on Two-Electron Processes 5

Current, i (μA)

IV III 0

–5

I

II –2.75

–2.5

–2.25

(a)

Current, i (μA)

IV III 0

I II

–5 –2.75

(b)

–2.5

–2.25

5

Current, i (μA)

II 0

–5 I

(c)

–2.7

–2.6

–2.5

–2.4

–2.3

–2.2

–2.1

FIgURE 11.9 Cyclic voltammograms at 50 mV s−1 with 0.24 mM 8: (a) R = Me, (b) R = Et, (c) R = iPr in THF/0.2 [NBu4][PF6] at a GC electrode: circles, experiment data; lines, digital simulation; potentials in V vs. Fc/Fc+. (From Dümmling, S. et al., Acta Chem. Scand., 53, 876, 1999. With permission.) R N N R 8

S S

2. Combined Structural and Medium Effects a. Reduction of [M2Mo(η5−C5Me5)2(S2C6H4)2(CO)2] (M = Co, Rh) The cobalt and rhodium Mo–bridged bimetallic compound 9 exhibits two-electron reductive behavior in which both structure change and medium effects are involved in determining the ordering of the ΔE1/2 values [33]. As indicated in the calculation-based structures of Scheme 11.6, the two semibridging

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Organic Electrochemistry

S S

–

–

–

+e

82–

–

–e S 5 4

2

–

3

S

+e–

1

–

–e–

S

S

8

8

–

+e – –e – R

–

N

S

–

N

S

–

R 82–

SCHEME 11.5 Structural changes during stepwise two-electron oxidation of the carbene-CS2 adduct. (From Dümmling, S. et al., Acta Chem. Scand., 53, 876, 1999. With permission.)

Mo S S S S Co Mo0 Co

+2e

–

–2e

CC OO

S

+C C + X X OO

Co C O

9,Co Two semibridging COs (a) + X

Co

S S S S Mo0 Co Co

–

9,Co2–.2Na+ Two bridging COs

H

Na

C

(b) : Bu4N+, Na+@18–crown–6 –

S S S S Co Mo0 Co CC OO

9,Co Two semibridging COs

+e–

S S S S 0 Co Mo Co

–e–

C C O O + X

9,Co– One semibridging CO One bridging CO

2–

S S SS Mo0 Co

+e–

Co

–e–

C C OO +

X

+

X

9,Co2– Two bridging COs

(c)

SCHEME 11.6 (a) Schematic diagram of the CO group coordination mode transition between 9,Co and 9,Co2− upon one-step 2e− redox reaction. (b) Optimized structure of 9,Co2−·2Na+ (triplet state) (C2 symmetry) in the CPK model. (c) Schematic diagram of the CO group coordination mode transitions between 9,Co, 9,Co−, and 9,Co2− upon two-step 1e− redox reactions. (From Muratsugu, S. et al., Chem. Sci., 2, 1960, 2011. With permission.)

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Influence of Molecular and Medium Effects on Two-Electron Processes

M

S S S S Mo M CC OO

9

(c)

(d)

(e)

ΔEp

0

Current (μA)

4 ΔE 0' ΔE 0' = 282 mV

–4 4

(f ) Current (μA)

ΔE 0' = 221 mV

–4 4

(g)

0 ΔE

–4 –8 4 0 –4 –8 10 5 0 –5 –10 –15 –2400

0' =

110 mV

ΔE 0' = –10 mV

(h)

ΔE 0' = 195 mV

ΔEpa

5 0 –5 –10 5

0

Current (μA)

(b)

E 10'

ΔEpc

ΔEp = 38 mV

0 –5 –10 5

ΔEp = 88 mV

0 –5 –10 –15

Current (μA)

(a)

Current (μA) Current (μA) Current (μA) Current (μA) Current (μA)

E 20'

M = Co M = Rh

ΔEp = 106 mV

0 –5 –10

ΔEp = 304 mV

(i) –2000 –1600 –1200 Potential (mV vs. Fc+/Fc)

–800

–2400

–2000 –1600 –1200 Potential (mV vs. Fc+/Fc)

FIgURE 11.10 Cyclic voltammograms of (a) 9,Co, in 0.1 M [NBu4][ClO4]/MeCN/toluene (1 : 1 v/v) at 253 K; (b) 9,Co, in 0.1 M [NBu4][ClO4]/THF at 258 K; (c) 9,Co, in 0.1 M [Na][BPh4]/MeCN/toluene (1 : 1 v/v) at 258 K; (d) 9,Co, in 0.1 M [Na][BPh4]/THF at 258 K; (e) 9,Co, in 0.1 M 18-crown-6/[Na][BPh4]/THF at 298 K; (f) 9, Rh in 0.1 M [Bu4][NClO4]–MeCN/toluene (1 : 1 v/v) at 253 K; (g) 9, Rh in 0.1M [Bu4N][ClO4]–THF at 258 K; (h) 9, Rh in 0.1M Na[BPh4]–MeCN/toluene (1 : 1 v/v) at 258 K; (i) 9, Rh in 0.1 M Na[BPh4]–THF at 258 K. (From Muratsugu, S. et al., Chem. Sci., 2, 1960, 2011. With permission.)

CO ligands of the neutral complex become traditional bridging carbonyls in the dianion. Whereas only one two-electron wave is seen under all conditions for the Rh complex, the Co complex may go from normal to inverted ordering if the medium is tuned. When the medium discourages ion pairing of the product anions with electrolyte cations (Figure 11.10, scans (a) and (b)), well-resolved one-electron processes are observed with ΔE1/2 values of 200 mV or more. This changes when sodium ions are introduced as the cation of the supporting electrolyte, Na[B(C6H5)4]. Ion-pairing stabilization of the dianion in tetrahydrofuran (THF) gave an inverted EE system (Figure 11.10, scan (d)), and the individual one-electron waves returned when 18-crown-6 was present, reducing the ion-pairing strength of Na+ (scan e). The reader is referred to the original article for discussion of the Rh analogue [33]. b. 1,1,1-(CO)3-2-Ph-closo-1,2,3,4-MnC3B7H9 (10) A case study of the joint effects of structure change and manipulations of the medium is found in the EE reduction of the tricarbadecaboranyl manganese compound

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1,1,1-(CO)3-2-Ph-closo-1,2,3,4-MnC3B7H9, 10 [34], for which structures of all three components of the E.T. sequence of Equation 11.24 are known. X-ray crystallography showed + e−

+ e−

1/ 2

1/ 2

(11.24)

2− − ⇀ ⇀ 10 ↽  10 ↽  10 E (1) E (2)

O O C C C Mn

O Ph

C

C C 10

that the carborane ligand is η6-coordinated to the metal center in the neutral compound but η4-coordinated in both the monoanion and dianion. Taking into account other ligand rearrangements, the structural changes in this series were viewed as gradual in going from 10 to 10− and finally 102− [34]. One might therefore expect roughly equivalent structural relaxation energies in the two E.T. processes, so that 10 would not fit into either of the two extremes treated in Section 2.F., which considered EE reactions having one dominant structure change. In such a case, explaining and manipulating the ΔE1/2 of the system comes down almost exclusively to effects of the electrolyte medium. Figure 11.11 reproduces some of the medium effects observed for this system. Beginning with scan (d), two separate 1e− processes are observed when the solvent is strongly donating (here, THF, [donor

ipc

ipc ipa

0

ipa

3 μA

–0.5

(a)

–1 E (V) vs. Fc

–1.5

–2

0

ipa

(c)

–1 E (V) vs. Fc

–1.5

–1

–1.5

–2

E (V) vs. Fc

ipc

1 μA

–0.5

–0.5

(b)

ipc

0

3 μA

ipa

–2

0 (d)

1.5 μA

–0.5

–1 E (V) vs. Fc

–1.5

–2

FIgURE 11.11 CV scans (0.1 V s−1) of 10 in different media: (a) CH2Cl2/[NEt4][B(C6H3(CF3)2)4], (b) CH2Cl2/ [NBu4][B(C6F5)4], (c) CH3CN/[NBu4][B(C6F5)4], (d) THF/[NBu4][B(C6F5)4]. (From Nafady, A. et  al., Organometallics, 26, 4471, 2007. With permission.)

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number, DN, = 20]) and the supporting electrolyte cation is weakly ion pairing (here, [NBu4]+). Both of these properties weaken the interaction with the dianion (relative to the monoanion), thermodynamically disfavoring the formation of the dianion and pushing E1/2(2) negative compared to E1/2(1). The ΔE1/2 value, which is 300 mV under these conditions, can be lowered by employing less donating solvents or more strongly ion-pairing electrolyte cations. Examples of the former are acetonitrile (DN = 14.1) [scan (c), ΔE1/2 = 110 mV] and dichloromethane (DN = 0) [scan (b), ΔE1/2 = 75 mV]. Replacing [NBu4]+ by [NEt4]+ in dichloromethane further lowers the ΔE1/2 value to 35 mV, and the CV scan (a) begins to take on the look of a single two-electron process. Medium-dependent tuning of ΔE1/2 in order to switch from a two-electron process to a one-electron process is of value from both electrochemical and synthetic viewpoints. The fact that, in the present case, the disproportionation constant for the monoanion 10 − goes from Kdisp = 8.3 × 10 −6 in THF/[NBu4][PF6] to Kdisp = 0.26 in CH 2Cl 2/[NBu4][B(C6F 3(CF 3)2) 4] aided design of the medium conditions that would facilitate preferential generation (and, ultimately, isolation) of either the monoanion 10 − or the dianion 10 2−. Thus, chemical reduction of 10 by Co(η5 −C5Me5)2 , when carried out in low-donor, low-polarity solvents (dichloromethane or toluene) gave the pure dianion 10 2−, whereas the same reducing agent in THF cleanly gave the monoanion 10 − [34]. 3. Dominant Medium Effects The practice of inducing changes in the ΔE1/2 values of two-electron systems by alterations of solvent and/or supporting electrolyte is widespread in electrochemical applications. The effects of altering the alkyl chain length in tetraalkylammonium ions or employing smaller, more strongly ion pairing, cations such as alkali metals has a long history in the study of anion-based electrochemistry [35]. However, an equivalent breadth of ion-pairing options for supporting electrolyte anions was only more recently made available with the introduction of weakly coordinating anion (WCA) based electrolytes, predominantly employing either [B(C6H3(CF3))4]− (11,  BArF24) or [B(C6F5)4]− (12, TFAB) [36]. Pertinent to the present subject is the influence of WCAs on the ΔE1/2 values of EE reactions involving cationic products. Owing to the fact that ion-pairing interactions of any electrolyte anion increase with increasingly positively charged electrode products, altering the ion pairing has a larger effect on the second oxidation than the first. The scale of possible ΔE1/2 changes was first demonstrated in the groundbreaking work of Mann et al. [36a], who determined that the Kdisp value for a trication of a dirhodium complex fell by a factor of over 107 in going from [PF6]− to [B(C6H3(CF3)2)4]− in CH2Cl2, corresponding to an effective increase in ΔE1/2 of over 400 mV. F F3C

F

F

F

F

CF3

F3C

F

CF3 B– CF3

F

F

B–

F

F3C F3C

F

F

F

F

F F

F

F

CF3 11

F

F F 12

a. Oxidation of Bis(fulvalene)dinickel (2) A more comprehensive study of the effects of both solvent and supporting electrolyte anion on anodic ΔE1/2 values was carried out on the electronically well-understood [37], single-site

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[Bu4N][Cl

[Bu4N][PF6]

[Bu4N][B(C6F5)4]

0.4

0

–0.4 V vs. Fc+/Fc

–0.8

FIgURE 11.12 CV scans at 0.1 V s−1 of bis(fulvalene)dinickel 2 in CH2Cl2, 0.1 M [NBu4]Cl, [NBu4][PF6], or [NBu4][B(C6F5)4]. (From Barrière, F. and Geiger, W.E., J. Am. Chem. Soc., 128, 3980, 2006. With permission.)

complex bis(fulvalene)dinickel, 2 [38]. As delineated in Equation 11.12, changes in ΔE1/2 may arise from changes in both the solvation and ion-pairing characteristics of the medium. In a selection of 45 combinations of nonaqueous solvents and supporting electrolytes [38], the ΔE1/2 values for the sequence 2/2 +/2 2+ were shown to go from a low of 212 mV (in anisole/[NBu4] Cl) to a high of 850 mV (in CH 2Cl 2 /Na[B(C6H 3(CF 3)2) 4]), an impressive change of 638 mV (or 14.6 kcal mol−1) in destabilization of the dication compared to the monocation. Figure 11.12 shows the effect of changing only the electrolyte anion in the medium. The dramatic increase in ΔE1/2 while going from Cl−, to [PF6]−, to [B(C6F5) 4]− (480 mV, Table 11.4) is ascribed to the weaker ion-pairing energies of the progressively larger, more charge-delocalized, electrolyte anions. It was noted that the traditional electrolyte anions ([PF6]−, [BF4]−, [ClO 4]−), often considered to be ion-pairing “innocent,” actually fall midway between halides and WCAs in ionpairing strength. These concepts have been reviewed [36(b)]. The ion-pairing effects that play a dominant role in the potential shifts in CH 2Cl 2 (ε = 8.9) are still present in THF (ε = 7.5), but the ΔE1/2 shifts are less (361 mV for the same three anions; see Table 11.4). This is readily explained by noting the increased donor strength of THF, which enhances the solvation effects compared to dichloromethane, thus facilitating the second oxidation process. Another informative set of measurements involved DMSO, in which there is very little change in ΔE1/2 with alterations of the electrolyte anion (only 56 mV for the three anions). In this case, the ion-pairing effects are minimized owing to the

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TAbLE 11.4 Changes in ΔE1/2 Values for the Successive One-Electron Oxidations of 2 in Different Solvents Solvent Dichloromethane Tetrahydrofuran Dimethylsulfoxide

Solvation

Ion Pairing

ΔΔE1/2 (mV)

ΔKdisp

Weak Strong Strong

Strong Strong Weak

480 361 56

1.3 × 108 1.3 × 106 8.9

Source: Data taken from Barrière, F. and Geiger, W.E., J. Am. Chem. Soc., 128, 3980, 2006. In each case, the spread of ΔE1/2 for electrolytes containing Cl−, [PF6]−, and [B(C6F5)4]− is given as ΔΔE1/2.

1000

Na BArF24 TBA TFAB

ΔE½ (mV)

800

TBA BArF24 TBA BPh4

600

TBA PF6 TBA ClO4

400

TBA triflate TBA BF4

200

TBA Br TBA Cl

0 CH2Cl2

THF

DMSO

FIgURE 11.13 Schematic representation of the available tuning window of ΔE1/2 values for complex 2 as a function of the supporting salt in dichloromethane, tetrahydrofuran, and dimethylformamide. (Reprinted from Barrière, F. and Geiger, W.E., J. Am. Chem. Soc., 128, 3980, 2006. With permission. Consult this reference for exact ΔE1/2 values.)

high dielectric constant of DMSO (ε = 47.2). Figure 11.13 is a pictorial representation of the influence of solvent on the ΔE1/2 values in solutions of different anions. Remarkably, the anion-based differences in Κdisp, which are 1.3 × 10 8 in CH 2Cl 2 , fall by over seven orders of magnitude in DMSO. 4. Effect of Slow E.T. Kinetics: Reduction of bis(hexamethylbenzene)ruthenium2+ As discussed in Section III and exemplified by the reduction of cyclooctatetraene (Section II.F.3), two-electron processes having inverted or compressed ΔE1/2 values frequently display at least one slow E.T. reaction owing to redox-induced structural changes. An organometallic example is the reduction of the bis(hexamethy1benzene)ruthenium dication, [(η 6 −C6Me 6)2Ru]2+, 13 2+, which has the idealized structures shown below as it is reduced first to 13 + and then to 13 [39]. The defining structural feature is the bending of the η 6 -coordinated arene ring of the monocation into the η4 -coordinated structure of the neutral complex. This structural rearrangement provides the relaxation energy to compress the ΔE1/2 value and also is the origin of the increased inner-sphere reorganization energy required for the second E.T. The combination of structural changes and inherent electronic effects [40] leads to slow heterogeneous E.T. kinetics for the 13 +/13 couple. Thus, the η 6/η4 structure change eases the thermodynamics but impedes the kinetics of formation of 13.

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+

+e–

+e– Ru

Ru

Ru

132+

13+

13

In dichloromethane, the two one-electron processes are thermodynamically resolved (ΔE1/2 ≈ 0.14 V) in slow to medium sweep rate CV scans (Figure 11.14), which also give evidence of the slow second charge-transfer step. The stronger solvation of cations by acetonitrile lowers ΔE1/2 to the inverted value of −0.03V and changes the CV scan dramatically, giving a single two-electron wave. At high sweep rates, however, the electrochemically irreversible second wave partitions from the first wave (Figure 11.15), giving a rare example of “kinetic differentiation” of the two E.T. reactions of an EE process [39].

B.

INVOLVING TWO IDENTICAL REDOX SITES

A special and important case of EE processes is that involving one-electron transfer of molecularly linked single redox sites. The structures shown below are readily recognized as precursors to mixedvalent systems that would be obtained by one-electron reduction of 14 or by one-electron oxidation of 15 or 16. These and a wide variety of loosely related compounds, all of which are inherently twoelectron systems, have been intensely studied since the original reports on the formally mixed-valent species [Ru2(NH3)10(μ-pyrazine)]5+ (the “Creutz/Taube Ion”) [41] and [biferrocene][ClO4], 15+ [42]. The spectroscopic and chemical behaviors of these molecules are tied closely to the degree of electronic communication between the redox centers, as reviewed elsewhere [43].

6+ Fe (H3N)5Ru N

N Ru(NH3)5 14

N

N

Fe 16 15

1. Electronic Communication through the Linkage Our discussion of the ΔE1/2 values of this family, modeled as 17 in Scheme 11.7, is conceptually simpler than the single-site redox case owing to the rarity of significant structural relaxation in the E.T. processes. This allows the emphasis to be on electronic and medium effects alone. In terms of electronic effects, a purely electrostatic model would suggest that the differences in ΔE1/2 values for the two redox sites will decrease as the length of the linker increases and as its ability to transmit charge, either through bonds or through space, decreases. (Note: Even if the site-to-site electronic interaction is completely eliminated, E1/2(1) and E1/2(2) will differ owing to a statistical effect, with ΔE1/2 = (RT/F) ln 2 (= 35 mV at 298 K) [16].) Until fairly recently, ΔE1/2 values greater than about 400 mV were taken to indicate that the two E.T. sites were sufficiently strongly interacting that the mixed-valent intermediate could be characterized as having a delocalized electronic structure (delocalized 17+ in Scheme 11.7). More sophisticated

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(a)

(d) 5×i

(b)

(e)

(f )

(c)

–0.5

–1.0

–1.5

V vs. Fc

–0.5

–1.0

–1.5

V vs. Fc

FIgURE 11.14 0.5 mM [132+][BF4] in dichloromethane: (a) 0.1, (b) 0.2, (c) 0.5, (d) 1.0, (e) 2.0, and (f) 5.0 V s−1. (From Pierce, D.T. and Geiger, W.E., J. Am. Chem. Soc., 114, 6063, 1992. With permission.)

analyses now treat the ΔE1/2 value simply as one part of the body of physical and spectroscopic information about the degree of site-to-site interactions in the compounds [43]. Unquestionably, however, ΔE1/2 values play an important part in defining the analytical and synthetic properties of these systems owing to radical disproportionation effects, as discussed in Section II. Here, we consider two factors that may often be manipulated to achieve desired ΔE1/2 values for two-site redox systems: the length of the molecular linker (or bridge) and the makeup of the electrolyte medium.

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(a)

(d) 20×i

(b)

(e)

(c)

(f )

–0.5

–1.0

–1.5 V vs. Fc

–0.5

–1.0

–1.5V vs. Fc

FIgURE 11.15 1 mM [132+][BF4] in CH3CN: (a) 0.1, (b) 0.2, (c) 0.5, (d) 1.0, (e) 2.0, and (f) 5.0 V s−1. (From Pierce, D.T. and Geiger, W.E., J. Am. Chem. Soc., 114, 6063, 1992. With permission.) 1+

LM 17

SCHEME 11.7

d LM lize oca l e D –e– ML –e– Lo cal ize LM d

0.5+

0.5+

ML 2+ –e–

LM

1+

1+

ML

1+ –e– 172+ 1+

0

ML

Delocalized and localized charges in mixed-valent members of two-electron process.

A glance at structures 14–16 suggests the broad structural variability that has, in fact, been utilized in molecular linkers. For an example of how the structure of the linker might affect ΔE1/2 values, consider the differences between biferrocene (15) and structures in which the two ferrocenyl moieties are connected by either sp3- or sp-carbon-based linkages. In going from 15 to bis(ferrocenyl)acetylene (18) and then to bis(ferrocenyl)ethane (3), the ΔE1/2 values measured in traditional electrolyte solutions decrease progressively from 330 [44] to 130 mV [42] to approximately 40 mV [45], in concert with the decreasing electronic communication expected with the changes in the bridge between the cyclopentadienyl ligands. Although these systems retain normal potential ordering with positive ΔE1/2 values, the dramatically higher values of Kdisp for 18+ and 3+ mean that

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Influence of Molecular and Medium Effects on Two-Electron Processes

nominally pure solutions of these radical cations will contain significant amounts of their neutral and dicationic E.T. partners, of particular importance in characterizing the chemical and physical properties of these mixed-valent species [43]. Inverted potential ordering is not normally observed for these types of systems, unless the linker itself undergoes E.T.-induced structural changes.

Fe

Fe

18

It is interesting to contrast the aforementioned case with one in which the carbon-based linkages are bonded directly to the metal centers of the terminal redox centers, rather than to the Cp rings. Here, we note the results of two related studies in which the lengths of alkynyl linkages directly connected to the metal of a half-sandwich redox center have been systematically altered. With either rhenium-based or iron-based redox centers (structures 19 and 20, respectively), large positive ΔE1/2 values are observed even when the alkynyl linker is eight carbons long, indicative of strong electronic communication between the metal centers (Scheme 11.8) [46,47]. The electrochemistry of these EE systems is relatively straightforward owing to their large potential spacings. We will now turn attention to systems with smaller potential spacings, and the ways in which the electrolyte medium can be manipulated to achieve chemically desired ΔE1/2 values. 2. Medium Effects a. Multi(ferrocenyl) Complexes Given the positive charge of the electrode products, ΔE1/2 values for the oxidation of multi(ferrocenyl) complexes can be expected to be quite sensitive to the ion-pairing strength of the electrolyte anion. Indeed, replacing a traditional anion such as [PF6]−, [BF4]−, or [ClO4]− with the WCA [B(C6F5)4]− increases the ΔE1/2 value significantly, often by hundreds of mV for bis(ferrocenyl) complexes Me5

Me5

Me5 Re

C

Me5 Fe

Re

C n

PPh3 NO

NO

P

PPh3

Fe

C

C

n P

P P

20

19 Compound 19 20

n 2 2

ΔE1/2 (V) 0.53 0.70

19 20

4 4

0.38 0.53

19 20

6 6

0.28 0.43

SCHEME 11.8 ΔE1/2 values for successive one-electron oxidations of alkynyl-linked organometallic redox centers, data taken from Reference 46 (compound 19) and Reference 47 (compound 20), for which P—P represents bis(diphenylphosphino)ethane.

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TAbLE 11.5 Representative Supporting Electrolyte Anion Effects on Differences in ΔE1/2 Values for bis(ferrocenyl) Compounds Compound

ΔE1/2 with Traditional Anion (mV)

Bis(fulvalene)diiron Biferrocene, 15 Bis(ferrocenyl)ethane, 3

590 330 ≈40

ΔE1/2 with [b(C6F5)4]− (mV) 970 530 180

“Traditional anion” refers to [PF6]−, [BF4]−, or [ClO4]−, with the solvent being either dichloromethane, acetonitrile, or mixed dichloromethane/acetonitrile. With [B(C6F5)4]–, the solvent was dichloromethane. See Reference 48 for details and original literature references.

(see Table 11.5) [48]. The medium-dependent sensitivity of ΔE1/2 values has been noted as a factor to be cautiously considered when evaluating potentially mixed-valent systems [38,49,50]. Let us use the general principles developed in this chapter to consider a potentially three-electron system terferrocene, 22.

Fe

Fe

Fe

22

b. Oxidation of Terferrocene, 22 Owing to the fact that the three ferrocenyl moieties in 22 are electronically coupled, three successive one-electron oxidations are observed for the four-member EEE E.T. series 22/22+/222+/223+. Figure 11.16 gives representative examples of how alterations in the medium affect the voltammetry of this system. A key change is that shown on the right side of the figure, where dichloromethane is the solvent and the electrolyte anion is either [PF6]− (c) or [B(C6F5)4]− (d), the overall ΔE1/2 for all three oxidations being altered by about 500 mV [48]. c. Fulvalene Dirhodium Complex 23 A final example taken from the chemistry of identical redox sites is that of the anodic EE system of (fulvalendiyl)bis(cyclooctadiene)dirhodium, 23. In this case, changes in the medium enabled fine tuning between compressed and inverted ΔE1/2 values [51].

Rh

Rh

23

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Influence of Molecular and Medium Effects on Two-Electron Processes

I (A)

1×10–5 0

–1×10–5 (a)

(c)

0

0

I (A)

1×10–5

I (A)

1×10–5

–1×10–5

–1×10–5 1.2

0.8

(b)

0.4 0 E (V vs. Fc/Fc+)

–0.4

1.2 (d)

0.8

0.4 0 E (V vs. Fc/Fc+)

–0.4

FIgURE 11.16 CV scans at 0.2 V s−1 for terferrocene (22) in different nonaqueous solutions: (a) reproduction of scan from Brown, G.M. et al., Inorg. Chem. 14, 506, 1975, in 1:1 CH3CN:CH2Cl2/0.1 M [NBu4][PF6]; (b) 0.4 mM 22 in CH3CN/0.1 M [NBu4][PF6]; (c) 0.4 mM 22 in CH2Cl2/0.1 M [NBu4][PF6]; (d) 0.4 mM 22 in CH2Cl2/0.1 M [NBu4][B(C6F5)4]. (From Camire, N. et al., J. Organometal. Chem., 637–639, 823, 2001. With permission.)

Consider the SWVs of Figure 11.17, which begin with the maximum separation (ΔE1/2 = 280 mV) in dichloromethane/[NBu4][B(C6H3(CF3)2) 4], scan (a). From there, addition of the [PF6]− anion [scan (b)], and then increasing amounts of the donor solvent 1,2-dimethoxyethane [scans (d) through (f)], all shift E1/2(2) negative with respect to E1/2(1), collapsing the voltammetric curves into the picture of an apparent single two-electron process. In the limit of pure glyme/[NBu4][PF6], an inverted potential ordering exists with ΔE1/2 = −50 mV [51].

VI.

AN INTEgRATED APPROACH TO MEDIUM EFFECTS ON ΔE1/2

Fe

Fe S

S Ni

S

S

Fe

Fe

24

Our broad discussion of extrinsic effects on ΔE1/2 values aids our understanding of the voltammetric behavior of a molecule having both multistep reductions and multistep oxidations, beginning from a neutral complex. Tetrakis(ferrocenyl)nickel dithiolene (24) has two reversible one-electron reductions located primarily on the nickel dithiolene center and four reversible one-electron oxidations located at the modestly interacting ferrocenyl moieties [38]. Systematic manipulation of the

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Organic Electrochemistry

0.5 μA/M

0.5 μA/M

0.4

0.2

(a)

0

–0.2

–0.4

–0.6

E (V) vs. Fc

0.1

0

(b)

–0.1

–0.2

–0.3

–0.4

0.6 μA/M

0.5 μA/M

0.1

0

(c)

–0.1

–0.2

–0.3

–0.4

–0.5

E (V) vs. Fc

0.1

0

(d)

–0.1

–0.2

–0.3

–0.4

0

–0.1

–0.2

–0.3

E (V) vs. Fc

–0.4

–0.5

–0.5

E (V) vs. Fc

0.7 μA/M

(e)

–0.5

E (V) vs. Fc

1 μA/M

0 (f )

–0.1

–0.2

–0.3

–0.4

–0.5

E (V) vs. Fc

FIgURE 11.17 Square-wave voltammograms (frequency 10  Hz, pulse height 25 mV) of 0.6  mM 23 at 1 mm GC disk in (a) CH2Cl2/0.1 M [NBu4][TFAB], (b) CH2Cl2/0.1 M [NBu4][TFAB] + 250 equiv of [NBu4] [PF6], (c) solvent mixture CH2Cl2/glyme (90: 10% v/v) + 0.1 M [NBu4][TFAB] and 250 equiv of [NBu4][PF6], (d) CH2Cl2/glyme (80%:20% v/v) + 0.1 M [NBu4][TFAB] and 250 equiv of [NBu4][PF6], (e) CH2Cl2/glyme (60%:40% v/v) + 0.1 M [NBu4][TFAB] and 250 equiv of [NBu4][PF6], (f) CH2Cl2/glyme (50%:50% v/v) + 0.1 M [NBu4][TFAB] and 250 equiv of [NBu4][PF6]. Glyme = 1,2-dimethoxyethane. (From Nafady, N. et al., Organometallics, 25, 1654, 2006. With permission.)

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Influence of Molecular and Medium Effects on Two-Electron Processes

[Bu4N][PF6]

V vs. Fc+/Fc 1

0.5

0

–0.5

–1

–1.5

–2

Na[B(C6H3(CF3)2)4]

FIgURE 11.18 CV scans at 0.1 V s−1 of tetraferrocenyl(nickel dithiolene) (24) in CH2Cl2, 0.1 M [Bu4N][PF6] and 0.02 M Na[B{C6H3(CF3)2}4] over the full potential range (oxidation and reduction). (From Barrière, F. and Geiger, W.E., J. Am. Chem. Soc., 128, 3980, 2006. With permission.)

solvent/supporting electrolyte medium allows either the maximizing or minimizing of ΔE1/2 values for both reductions and oxidations of 24. In terms of solvent effects, strong donor solvents (e.g., ethers) favor increased separation of cathodic ΔE1/2 values but compression of anodic ΔE1/2 values, whereas the opposite effect accrues for weak donor/strong acceptor solvents such as halocarbons. Ion-pairing interactions between the redox products and the supporting electrolyte ions also play an important role in tuning the ΔE1/2 values, providing that the experiments are carried out in solvents of modest to low dielectric constant (ε ≈ 10 or less). Examples are given in Figure 11.18. Although dichloromethane is a favorable solvent for the separation of oxidation processes, the four anodic waves arising from the E.T. sequence 24/24+/242+/243+/244+ are unresolved in CH2Cl2/ [NBu4][PF6] (upper left in figure) owing to relatively strong ion pairing of the product cations with [PF6]−. The two cathodic waves arising from 24/24 −/242− are, however, well resolved (ΔE1/2 = 0.78 V, upper right) owing to weak ion-pairing interactions of the product anions with [NBu4]+. The situation is reversed when the electrolyte salt is Na[B(C6H3(CF3)2]4 owing to the weak ion-pairing ability of [B(C6H3(CF3)2]4− with 24 n+ and the strong ion pairing of Na+ with 24 n− (lower scans of Figure 11.18). Interestingly, the reductions go from an obviously separate EE process to a single two-electron wave with the introduction of a sodium ion-based electrolyte. Weighing the quantitative details of these effects allowed formulation of a set of rules for the medium in increasing potential separations (i.e., enlarging ΔE1/2 values) of either oxidations or reductions. For product cations, employ (1) a lower polarity solvent of low donor number, (2) a weakly ion-pairing electrolyte anion, and (3) a strongly ion-pairing electrolyte counter cation. For product anions, employ (1) a lower polarity solvent of low acceptor number, (2) a weakly ion-pairing electrolyte cation, and (3) a strongly ion-pairing electrolyte counter anion. Given the inverse complementarity of these effects, the authors proposed a “mirror image” model for the effects of solvent and supporting electrolyte on ΔE1/2 values of both reductions and oxidations (see Figure 11.19) [38].

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Generation of cations

Generation of anions Intermediate cations in electrolyte + for example: Me4N

Intermediate anions in electrolyte – for example: PF6

ΔE1/2

0V

Low

Large anion in electrolyte – for example: TFAB

Large cations in electrolyte + for example: Bu4N

ΔE1/2

Small anions Small cations in electrolyte in electrolyte for example: Cl– for example: Na+

High

High

0V

Low

Solvent polarity Zone for conversion of sequential one-electron waves into a single two-electron wave.

FIgURE 11.19 Mirror image model of electrolyte effects on ΔE1/2. (From Barrière, F. and Geiger, W.E., J. Am. Chem. Soc., 128, 3980, 2006. With permission.)

ACKNOWLEDgMENTS The authors are grateful to the National Science Foundation for support during the writing of this chapter.

REFERENCES 1. Evans, D.H. Chem. Rev. 1990, 90, 739–751. 2. (a) Zusman, L.D.; Beratan, D.N. J. Phys. Chem. A 1997, 101, 4136–4141; (b) Lambert, C. Chem. Phys. Chem. 2003, 4, 877–880; (c) Gileadi, E. J. Electroanal. Chem. 2002, 532, 181–189. 3. Evans, D.H. Chem. Rev. 2008, 108, 2113–2144. 4. Evans, D.H.; Hu, K. J. Chem. Soc., Faraday Trans. 1996, 92, 3983–3990. 5. Hill, M.G.; Rosenhein, L.D.; Mann, K.R.; Mu, X.H.; Schultz, F.A. Inorg. Chem. 1992, 31, 4108–4111. 6. Fernandez, J.B.; Zhang, L.Q.; Schultz, F.A. J. Electroanal. Chem. 1991, 297, 145–161. 7. For a discussion of the Born equation as applied to solvation energies see Bockris, J.O.’M.; Reddy, A.K.N. In: Modern Electrochemistry: Ionics. Plenum Press, New York, 1998, Vol. 1, pp. 204–207. 8. Fry, A. Electrochem. Commun. 2005, 7, 602–606. 9. Jensen, B.S.; Parker, V.D. J. Am. Chem. Soc. 1975, 97, 5211–5217. 10. Evans, D.H.; O’Connell, K.M. In: Electroanalytical Chemistry. A Series of Advances, Bard, A.J. (Ed.), Marcel Dekker, New York, 1986, Vol. 14, pp. 113–207. 11. For early studies see: (a) Katz, T.J. J. Am. Chem. Soc. 1960, 82, 3784–3785; (b) Allendoerfer, R.D.; Rieger, P.H. J. Am. Chem. Soc. 1965, 87, 2336–2344; (c) Paquette, L.A.; Wright III, D.; Traynor, S.G.; Taggart, D.L.; Ewing, G.D. Tetrahedron 1976, 32, 1885–1891. 12. Huebert, B.J.; Smith, D.E. J. Electroanal. Chem. 1971, 31, 333–348. 13. Smith, W.H.; Bard, A.J. J. Electroanal. Chem. 1977, 76, 19–26. 14. Baik, M.-H.; Schauer, C.K.; Ziegler, T. J. Am. Chem. Soc. 2002, 124, 11167–11181. 15. Evans, D.H. Acta Chem. Scand. 1998, 52, 194–197.

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Influence of Molecular and Medium Effects on Two-Electron Processes 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43.

44. 45. 46. 47. 48. 49. 50. 51. 52. 53.

431

Flanagan, J.B.; Margel, S.; Bard, A.J.; Anson, F.C. J. Am. Chem. Soc. 1978, 100, 4248–4253. Sun, H.; Chen, W.; Kaifer, A.E. Organometallics 2006, 25, 1828–1830. Li, D.; Keresztes, I.; Hopson, R.; Willard, P.G. Acc. Chem. Res. 2009, 42, 270–280. Nicholson, R.S.; Shain, I. Anal. Chem. 1964, 36, 706–723. Bard, A.J.; Faulkner, L.R. Electrochemical Methods, John Wiley & Sons, New York, 2nd Edn., 2001, Chapter 6, p. 226ff. Polcyn, D.S.; Shain, I. Anal. Chem. 1966, 38, 370–375. Richardson, D.E.; Taube, H. Inorg. Chem. 1981, 20, 1278–1285. Bard, A.J.; Santhanam, K.S.V. In: Electroanalytical Chemistry. A Series of Advances, Bard, A.J. (Ed.), Marcel Dekker, New York, 1970, Vol. 4, pp. 215–315. Lingane, P.J. Anal. Chem. 1964, 36, 1723–1726. Amatore, C.; Azzabi, M.; Calas, P.; Jutand, A.; Lefrou, C.; Rollin, Y. J. Electroanal. Chem. 1990, 288, 45–63. Koelle, U.; Fuss, B.; Rajasekharan, M.V.; Ramakrishna, B.L.; Ammeter, J.H.; Boehm, M.C. J. Am. Chem. Soc. 1984, 106, 4152–4160. (a) Geiger, W.E. Acc. Chem. Res. 1995, 28, 351–357 (b) Bowyer, W.J.; Geiger, W.E. J. Electroanal. Chem. 1988, 239, 253–271. Kraiya, C.; Evans, D.H. J. Electroanal. Chem. 2004, 565, 29–35. Evans, D.H.; Busch, R.W. J. Am. Chem. Soc. 1982, 104, 5057–5062. Yoshida, Z.; Kawase, T.; Awaji, H.; Sugimoto, I.; Sugimoto, T.; Yoneda, S. Tetrahedron Lett. 1983, 24, 3469–3472. Bellec, N.; Boubekeur, K.; Carlier, R.; Hapiot, P.; Lorcy, D.; Tallec, A. J. Phys. Chem. A. 2000, 104, 9750–9759. Dümmling, S.; Speiser, B.; Kuhn, N.; Weyers, G. Acta Chem. Scand. 1999, 53, 876–886. Muratsugu, S.; Sodeyama, K.; Kitamura, F.; Tsukada, S.; Tada, M.; Tsuneyuki, S.; Nishihara, H. Chem. Sci. 2011, 2, 1960–1968. Nafady, A.; Butterick III, R.; Calhorda, M.J.; Carroll, P.J.; Chong, D.; Geiger, W.E.; Sneddon, L.G. Organometallics 2007, 26, 4471–4482. Heinze, J. In: Encylopedia of Electrochemistry, Bard, A.J.; Stratmann, M. (Eds.), Wiley-VCH Publishers, Weinheim, Germany, 2004, Vol. 8 (Schäfer, H.J., Ed.), p. 93ff. (a) Hill, M.G.; Lamanna, W.M.; Mann, K.R. Inorg. Chem. 1991, 30, 4687–4690 (b) Geiger, W.E.; Barrière, F. Acc. Chem. Res. 2010, 43, 1030–1039. Smart, J.C.; Pinsky, B.L. J. Am. Chem. Soc. 1977, 99, 956–957. Barrière, F.; Geiger, W.E. J. Am. Chem. Soc. 2006, 128, 3980–3989. Pierce, D.T.; Geiger, W.E. J. Am. Chem. Soc. 1992, 114, 6063–6073. Lord, R.L.; Schauer, C.K.; Schultz, F.A.; Baik, M.-H. J. Am. Chem. Soc. 2011, 133, 18234–18242. Creutz, C.; Taube, H. J. Am. Chem. Soc. 1969, 91, 3988–3989. LeVanda, C.; Cowan, D.O.; Bechgaard, K. J. Am. Chem. Soc. 1975, 97, 1980–1981. (a) Creutz, C. In: Progress in Inorganic Chemistry, Lippard, S.J. (Ed.), John Wiley & Sons, New York, 1983, Vol. 30, p. 1ff; (b) Barlow, S.; O’Hare, D. Chem. Rev. 1997, 97, 637–669; (c) Chen, P.; Meyer, T.J. Chem. Rev. 1998, 98, 1439–1477; (d) Demadis, K.D.; Hartshorn, C.M.; Meyer, T.J. Chem. Rev. 2001, 101, 2655–2686; (e) Nelsen, S.F. Chem. Eur. J. 2000, 6, 581–588. Morrison, Jr., W.H.; Krogsrud, S.; Hendrickson, D.N. Inorg. Chem. 1973, 12, 1988–2004. This ΔE1/2 was reported as 350 mV in Reference 42. For reviews of linked metallocene-based redox centers, see Reference 43b and Nishihara, H. Adv. Inorg. Chem. 2002, 53, 41–86. Meyer, W.E.; Amoroso, A.J.; Horn, C.R.; Jaeger, M.; Gladysz, J.A. Organometallics 2001, 20, 1115–1127. Paul, F.; Lapinte, C. Coord. Chem. Rev. 1998, 178–180, 431–509. Camire, N.; Mueller-Westerhoff, U.T.; Geiger, W.E. J. Organometal. Chem. 2001, 637–639, 823–826. Barrière, F.; Camire, N.; Geiger, W.E.; Mueller-Westerhoff, U.T.; Sanders, R. J. Am. Chem. Soc. 2002, 124, 7262–7263. D’Alessandro, D.M.; Keene, F.R. J. Chem. Soc., Dalton Trans. 2004, 3950–3954. Nafady, N.; Chin, T.T.; Geiger, W.E. Organometallics 2006, 25, 1654–1663. Brown, G.M.; Meyer, T.J.; Cowan, D.O.; LeVanda, C.; Kaufman, F.; Roling, P.V.; Rausch, M.D. Inorg. Chem. 1975, 14, 506. Nicholson, R.S. Anal. Chem. 1965, 37, 1351.

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12

Electrochemically Driven Supramolecular Devices Paola Ceroni, Alberto Credi, and Margherita Venturi

CONTENTS I. II.

Introduction .......................................................................................................................... 433 Electrochemical Analysis of Supramolecular Systems ........................................................ 434 A. Intermolecular Interactions ........................................................................................... 435 B. Molecular Encapsulation of Electroactive Units .......................................................... 435 C. Redox-Controlled Supramolecular Switching .............................................................. 435 D. Electroactive Species on a Solid Support ..................................................................... 436 III. Host–Guest Systems ............................................................................................................. 437 A. Hydrogen-Bonding Interactions .................................................................................... 437 B. Metal Ion Coordination................................................................................................. 438 1. Metal Ion Translocation ......................................................................................... 439 2. Anion Translocation...............................................................................................440 3. Helicate Assembly and Disassembly ..................................................................... 441 C. Hydrophobic Interactions .............................................................................................. 442 D. Charge-Transfer Interactions ........................................................................................444 E. Host-guest Systems Working on Surfaces ....................................................................449 IV. Molecular Machines ............................................................................................................. 450 A. Basic Concepts .............................................................................................................. 450 B. Systems Based on Rotaxanes ........................................................................................ 450 1. Molecular Shuttles ................................................................................................. 451 2. Ring Pirouetting Motion ........................................................................................ 456 C. Systems Based on Catenanes ........................................................................................ 456 D. Molecular Machines Working on Surfaces ..................................................................466 1. Molecular Machines Immobilized on Electrodes..................................................466 2. Solid-State Electronic Circuits ..............................................................................468 3. Electrochemically Induced Shuttling in Single Rotaxane Molecules ................... 471 4. Electrically Driven Directional Motion of a Single Molecular Machine .............. 472 V. Concluding Remarks ............................................................................................................ 473 Acknowledgments.......................................................................................................................... 474 References ...................................................................................................................................... 474

I. INTRODUCTION Supramolecular chemistry, according to its most popular definition, is “the chemistry beyond the molecule, bearing on organized entities of higher complexity that result from the association of two or more chemical species held together by intermolecular forces” [1]. The field has developed at an astonishingly fast rate during the last three decades, and it soon became evident that a definition strictly based on the nature of the bond that links the components would be limiting.

433

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Many scientists, therefore, started to distinguish between what is molecular and what is supramolecular based on the degree of intercomponent interactions [2,3]. In a general sense, one can say that with supramolecular chemistry, there has been a shift in focus from molecules to molecular assemblies or multicomponent structures driven by the emerging of new functions. In the frame of research on supramolecular systems, the idea began to arise in a few laboratories [4–6] that the concepts of device and machine could be applied at the molecular level [7]. In other words, molecules might be used as building blocks for the assembly of multicomponent structures exhibiting novel and complex functions that arise from the cooperation of simpler functions performed by each component. This strategy, encouraged by a better understanding of biomolecular devices, has been implemented on a wide variety of chemical systems, leading to highly interesting results [8]. As a matter of fact, the molecular bottom-up construction of nanoscale devices and machines has become one of the most stimulating challenges of nanoscience. Such achievements have been made possible because of the substantial progresses obtained in other areas of chemistry and physics—particularly concerning the synthesis and characterization of complex chemical systems and the study of surfaces and interfaces. In this perspective, electrochemistry is a very powerful tool not only for characterizing a supramolecular system but also for the device operation. Indeed, molecular devices, as their macroscopic counterparts, need energy to operate and signals to communicate with the operator. Electrochemistry is an interesting answer to this dual requirement: it can be used to supply the energy needed to make the system work, and by means of the various electrochemical techniques (e.g., voltammetry), it can also be used to read the state of the system, controlling and monitoring the operation performed by the device. Furthermore, electrodes represent one of the best ways to interface molecular-level systems to the macroscopic world, a feature which is important for future applications. Hence, it is not surprising that the marriage of electrochemistry and supramolecular chemistry has produced a wealth of very interesting devices and functions, thereby generating new scientific knowledge and raising expectations for practical applications in energy conversion, information and communication technologies, advanced materials, diagnostics, and medicine. Our aim with this chapter is to provide a picture of the potentialities and applications of electrochemical methods for both investigating and operating multicomponent chemical systems. We start with a brief discussion of the role and potentialities of electrochemistry in the study of supramolecular systems. We then illustrate examples taken from electrochemical research applied to supramolecular and nanoscale systems, with particular attention to properties and functions. Although these examples cover a wide range of topics, we do not even attempt to be comprehensive, and we apologize from the beginning with the colleagues who may feel that their work has been omitted. Furthermore, in some cases, the selected examples are dated, but we decided to privilege those characterized by a high educational value. Several books [9–14] and reviews [15–22] are available for more thorough discussions on the fundamentals of electrochemical methods and their application to supramolecular systems and materials.

II.

ELECTROCHEMICAL ANALySIS OF SUPRAMOLECULAR SySTEMS

The primary task of most electrochemical techniques is the determination of redox potentials, which are then correlated to intrinsic electronic properties of molecules such as the energy of HOMO and LUMO levels. Electrochemical measurements give also access to other important quantities: diffusion coefficients, interfacial properties, and thermodynamic and kinetic constants of reactions coupled with the redox process. The knowledge of these parameters is extremely useful for understanding the behavior of many kinds of supramolecular assemblies. In this section, we discuss a few important issues related to the role and potentialities of electrochemistry for the study of multicomponent molecular systems.

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Electrochemically Driven Supramolecular Devices

A.

435

INTERMOLECULAR INTERACTIONS

The intermolecular forces that are at the basis of supramolecular chemistry, encompassing dipole– dipole, ion–dipole, ion–ion, hydrogen bonding, π–π, and cation–π interactions, are largely electrostatic in character. Solvophobic (hydrophobic) and van der Waals effects can also be significant in self-assembly processes, especially those responsible for the formation of monolayers. First of all, one should be aware that the conditions usually required for electrochemical experiments (solvents with reasonably high dielectric constants, large concentration of supporting electrolyte) can significantly affect the extent of the intermolecular interactions. Moreover, the redox potential measured for a given species may differ from the true value due to interactions with the ions of the supporting electrolyte. Second, given the nature of the forces that govern supramolecular phenomena, a change in the oxidation state of an electroactive species can greatly influence its propensity to interact with other molecular components. Because of such an interplay, redox potential values can give valuable information on intermolecular interaction energies. On the other hand, electrochemically driven changes of the oxidation state of molecular components can be used to direct and control their supramolecular interactions. The dual utility of electrochemistry as a detector and as an effector tool for supramolecular systems will emerge clearly from the examples discussed in the next sections.

B.

MOLECULAR ENCAPSULATION OF ELECTROACTIVE UNITS

The general meaning of molecular encapsulation is the placement of a molecule inside a much larger one. Encapsulation can be obtained by taking advantage of supramolecular interactions: if the smaller component (referred to as the guest) is held inside the larger one (the host) for an experimentally significant period of time, one can say that the host encapsulates the guest. Another approach to molecular encapsulation is to covalently attach large substituents (usually, polymeric structures or dendritic branches) to the species of interest. In both cases, encapsulation may result in site isolation of the molecule, that is, the segregation or protection from the solvent and other species present in the medium. The encapsulation of redox-active species can have a profound influence on their electrochemical properties [17,21,23]. In general, both the thermodynamic (i.e., the redox potential values) and the kinetic (i.e., the heterogeneous rate constants) aspects of the electron transfer process involving the encapsulated species will be different from those characteristic of the free form. Another distinctive feature of molecular encapsulation that can be evidenced by electrochemical techniques is the decrease of the apparent diffusion coefficient of the redox-active species when it is surrounded by a large molecular load. Moreover, in the case of noncovalent encapsulation, the dynamic nature of the self-assembly equilibria is important in determining the electrochemical response of the system, as the electron transfer can potentially involve both the encapsulated and free guest.

C. REDOX-CONTROLLED SUPRAMOLECULAR SWITCHING The oxidized and reduced states of an electrochemically switchable guest (or host) molecule generally exhibit a different degree of affinity for a host (or guest). Hence, the oxidation state of the redox-active species influences the thermodynamic stability of its complex with the other molecular component, and an electrochemically induced change in the oxidation state will cause a supramolecular switching event. When the binding interaction is sufficiently strong, the electrochemical behavior may clearly reflect the presence of two redox couples; in other words, the bound and free species may exhibit different redox potentials. In order for the switching to be effective, the redox-controlled species should exhibit fast electrontransfer kinetics. Without reversible kinetics, the switching would become too slow to be useful. A second requirement is that at least one of the redox states must interact strongly with the molecular partner.

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Organic Electrochemistry KH∙G G

+

H∙G

H

+e–

+e– E °' f

–e–

–e– KH∙G–

G–

+

E c°'

H ∙ G–

H

FIgURE 12.1 Square scheme mechanism describing an electrochemically switchable host–guest system. Horizontal and vertical processes are the chemical equilibria and the electron transfer processes, respectively.

It is easy to show that these two conditions are fulfilled for any redox-switchable system, from receptor–ion complexes to molecular shuttles. An electrochemically switchable system is usually described with a thermodynamic square scheme (Figure 12.1) that represents two electron transfer processes coupled with two homogeneous chemical equilibria (electron transfer, chemical reaction, electron transfer, chemical reaction; ECEC mechanism). We can assume, for example, that the guest G is electroactive and its binding with the host H is switched from low to high upon reduction. Therefore, the stability constant KH·G− is larger than KH·G; the ratio KH·G−/KH·G is defined as the binding enhancement and its magnitude can be estimated from the difference in the formal potentials of the free Ef°′ and complexed Ec°′ guest equation as follows: °′

°′

F ( Ef − Ec ) − K H ⋅G − RT =e K H ⋅G

(12.1)

Both Ef°′ and Ec°′ values are usually approximated by the half-wave potentials; the larger the difference in the half-wave potentials, the greater the value of the binding enhancement. The magnitude of KH·G determines whether the free or complexed guest is reduced at the electrode surface. When KH·G is large, the already formed H·G complex is electrochemically switched to a higher affinity state, H·G−. In this case, the diffusion of the host species is not relevant. Conversely, if KH·G is small, the species undergoing reduction will be the free guest G, which will subsequently bind the host to yield H·G−. In this situation, after the reduction of G to the high binding state, the complexation process is essentially controlled by the diffusion of the host toward the reduced guest. In general, two separate voltammetric waves for the free and complexed guest species will not be observed for systems with a low KH·G. In this binding regime, it may be important to consider the value of the rate constants for complexation and decomplexation. The typical electrochemical response will consist of a shift of the half-wave potential for the free guest species as the host is added to the solution. A thorough discussion of these processes is beyond the scope of this chapter and can be found elsewhere [10,14,15].

D. ELECTROACTIVE SPECIES ON A SOLID SUPPORT The organization and deposition of molecules on surfaces are important both for fundamental studies and for the development of platforms to interface molecular systems with the macroscopic world, thus enabling the construction of practical devices. Self-assembled monolayers (SAMs), Langmuir monolayers, and Langmuir–Blodgett multilayer films are widely used techniques to immobilize molecules on surfaces or at interfaces [24,25]. Molecular recognition events involving systems in which either the host or the guest component is bound to the surface of a solid electrode can be probed by electrochemical techniques [26]. Moreover, in many cases, the assembly–disassembly process can also be controlled by modulating the potential applied to the electrode [10,14,15,18,27,28]. Systems of this kind are extremely attractive for the construction of responsive surfaces that can find applications in, for example, chemo(bio)sensing, materials science, and drug delivery.

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III. HOST–gUEST SySTEMS As already mentioned in Section II in supramolecular chemistry, the term host refers to the larger and more structurally complex of two binding partners, while the term guest refers to the  smaller, less complex binding partner. In this section, examples of host–guest systems in which the strength of the binding interaction can be modulated electrochemically (see Figure 12.1) are reported. A redox process can affect the host–guest interaction in two ways: either the redox couple is directly involved in the binding with the other substrate or the charge perturbation brought about by the redox process modifies the binding constant via an electrostatic effect or a change in the conformation of the host or the guest. While the former approach is likely to give a stronger redox perturbation on the binding constant, the latter enables the design of the binding site independently from that of the redox-active moiety. In the following discussion, host–guest systems are divided according to the nature of interaction occurring between the host and the guest, although in several cases, interactions of different types are responsible for the association.

A.

HYDROGEN-BONDING INTERACTIONS

An example in which the redox site is directly involved in the binding process is constituted by the host–guest system [1•2] [29]. Azoflavin 1 strongly interacts with host 2, which contains two guanidinium groups providing four H bonds for azoflavin 1, in addition to the three setup to bind the imide portion of azoflavin. Because of these strong interactions, the system is characterized by a high association constant: KH·G = 1.9 × 104 M−1 in 20% CH3CN/CH2Cl2. Reduction of the azoflavin to its radical anion increases the binding strength even further, giving a ΔE1/2 = 317 mV. This potential shift corresponds to a binding enhancement of 2.2 × 105 and a KH·G– = 4.3 × 109 M−1 (see Equation 12.1), certainly one of the largest binding constants ever reported for these types of H-bonded complexes. In a similar approach, flavin 3 can interact with the simple receptor 2,6-diamidopyridine by hydrogen bond formation. This receptor unit bearing an alkyl chain terminated by a thiol group can be placed at the surface of a gold nanoparticle (4) [30]. The host–guest interaction is modulated by reducing the flavin in the presence of the functionalized gold nanoparticles. From the electrochemical studies, the corresponding association constant KH·G − can be estimated: it is about 20-fold greater than KH·G (4500 M−1 vs 196 M−1). The redox modulation process is selective as the reduction of a methylated flavin analog (methylated at the N(3) imide position to disrupt hydrogen bonding) is relatively unaffected by the presence of the same functionalized gold nanoparticles [30]. H

H N

H H

N H

N

O

+N

H22C12 N

H

1

+

N

N

H N

H N

H

H

N

N

O

N

H H

H

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N

N

N

N 2

CH N 6 13 C6H13

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Organic Electrochemistry

An example in which the binding and redox sites are distinct is represented by the host–guest system formed by compounds 5 and 6 [31]. 3 N

N

O H N

N

O

N

H N

O H

N O

S S S

4

Host molecule 5 contains a tetrathiafulvalene (TTF) moiety that undergoes two reversible oxidations [32], directly linked to an imide derivative able to bind diaminopyridines like 6. R

O N

H O S

S

S

S

N H N

O

N H

5

N Bu

O

R 6

Oxidation of the TTF decreases the H bonding by removing electron density from the electrondonating imide carbonyls; indeed, a modest +30 mV E1/2 shift for the TTF0/+ couple is observed in NBu4PF6/CH2Cl2, indicating a threefold decrease in binding strength upon oxidation.

B.

METAL ION COORDINATION

The redox process of a host–guest system can not only weaken the interaction of the two molecular components but also cause a movement of the guest inside the host. Representative examples are based on metal ion coordination, and involve either the translocation of a metal ion (guest) from one coordinating site to another one within the host or the rearrangement of the ligand (host) around the metal ion, as in the case of helicate assembly and disassembly. The role of the redox-active unit is usually played by the metal, in particular a transition metal ion. The requirements are (1) a fast and reversible one-electron transfer process and (2) different electronic and/or geometrical preferences of the two oxidation states of the metal center in the complex formation. Because of these essential features, most examples rely on the Fe(III)/Fe(II) and Cu(II)/Cu(I) redox couples. As to the coordination preferences [33], Fe(III) is a high-spin d5 cation, which cannot take advantage from ligand field stabilization energy (LFSE) effects and is therefore inclined to establish mainly electrostatic interactions, while Fe(II) is a low-spin d6 cation, which profits at most from LFSE in an octahedral coordinative environment and, in the presence of π-acceptor ligands, can also exert back donation.

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Electrochemically Driven Supramolecular Devices

For the copper-based couple, the Cu(II) ion (electronic configuration d9) is ready to adapt itself to any coordination number (4, 5, 6) and geometry offered by the coordinating environment, in order to profit at most from LFSE. Conversely, the Cu(I) ion (d10) requires a tetrahedral geometry because of inter-ligand steric repulsions. 1. Metal Ion Translocation A typical example of ion translocation is represented by a redox-active metal ion that moves from a given coordination environment to another one of the same supramolecular system, as a result of an electrochemical input (Figure 12.2). A supramolecular system for metal ion translocation must contain two distinct binding compartments differing for their coordinating properties: for example, one is a soft receptor and the other is a hard one. The metal ion must possess two adjacent oxidation states of comparable stability: one of soft nature (Mn+) and the other with hard character (M(n+1)+). In a solution containing equimolecular amounts of the reduced metal ion Mn+ and of the two-compartment supramolecular system, the metal ion Mn+ will occupy the soft compartment. When Mn+ is oxidized, the hard ion M(n+1)+ will move to the nearby hard compartment to increase its stabilization. Consecutive oxidation and reduction processes would make the metal center shuttle back and forth along a defined route. The translocation process is described by the electrochemical square scheme shown in Figure 12.1. A Fe(III)/Fe(II)-driven translocation process is observed with ligand 7 (Figure 12.3) [34], which consists of a 4-methylphenol platform, to which two different terdentate coordinating units have been appended in ortho positions: one consists of a tertiary amine nitrogen atom and two phenolate oxygen atoms and the other possesses one tertiary amine nitrogen atom and two pyridine nitrogen atoms. Upon addition of one equivalent of Fe(ClO4)3 to a CH3CN solution of 7, in the presence of collidine that guarantees deprotonation of all the phenolic groups, the Fe(III) cation is coordinated by the left compartment by three oxygen atoms, one nitrogen atom, and two solvent molecules (S), Hard

Soft +e–

M(n+1)+

Mn+ –e–

FIgURE 12.2 Scheme of a host system containing a soft and a hard receptor compartment for binding metal ions Mn+ and M(n+1)+, respectively. Metal ion translocation takes place by a reversible metal-centered redox process.

O–

N 3+

– O O –

Hard

N

+e– N

S S N

–e–

–

N – O – O

O

N

S

N

2+ S

N

Soft 7a

7b

FIgURE 12.3 Translocation motion of an iron ion within a two-compartment ligand. The left compartment has a harder character compared to the right one, so that the former hosts Fe(III) ions, while the latter is suitable for Fe(II) ions. Solvent molecules (S = CH3CN) complete the coordination sphere of the metal ion. The central phenolate oxygen atom is shared by the two compartments and acts as a pivot in the translocation process.

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i(A)

Organic Electrochemistry

FeIII

–0.8

–0.4 E/V

0

FeII

FIgURE 12.4 Cyclic voltammogram of a CH3CN/NBu4ClO4 solution of 7 (4 × 10 −3 M), containing one equivalent of Fe(ClO4)2 and collidine. Working electrode: platinum disk (5 mm diameter); scan rate: 0.1 V s−1; reference electrode: AgNO3/Ag in CH3CN. (From Ceroni, P., Credi, A., Venturi, M., and Schalley, C.A. eds.: Analytical Methods in Supramolecular Chemistry. pp. 371–457. 2012. Copyright Wiley-VCH Verlag GmbH & Co. KGaA. Reproduced with permission.)

giving rise to a six-coordinated complex (7a in Figure 12.3). On the other hand, if a stoichiometric amount of Fe(ClO4)2 is added to a solution of 7 containing collidine, the Fe(II) center is coordinated by the right compartment with three nitrogen atoms, one oxygen atom of the phenolate group, and two solvent molecules (7b in Figure 12.3). The Fe(II) ion forms again a six-coordinated complex, but with softer ligands compared to the Fe(III) ion. Translocation was carried out electrochemically and was investigated through cyclic voltammetry experiments in a CH3CN solution containing 0.1 M NBu4ClO4. The cyclic voltammogram recorded at 0.1 V/s (Figure 12.4) shows an anodic and a cathodic peak, both exhibiting chemically irreversibility according to the square mechanism illustrated in Figure 12.1. At a potential of −0.8 V, the Fe(II) ion stays in the right compartment (7b in Figure 12.3). On increasing the potential, oxidation to Fe(III) takes place, with an anodic peak at ca. −0.2 V. No corresponding cathodic peak is observed upon reversing the potential scan because of the coupled structural rearrangement, that is, the movement of the Fe(III) ion to the left compartment (7a in Figure 12.3). The corresponding Fe(III)/Fe(II) reduction is made more difficult and takes place at a more negative potential (Epc ca. −0.7 V). After this point, Fe(II) moves back to its original compartment and the system is ready for a further cycle. CV studies at high scan rates indicated that the lifetime of the translocation is shorter than 10 ms. The high translocation rate can be ascribed to the beneficial assistance of the central phenolate oxygen atom, which keeps the iron center coordinated over the course of the direct and reverse motion, thus reducing the energy of the transition state. 2. Anion Translocation Not only metal cations but also anions can be translocated within ditopic systems, taking advantage of the coordination to a redox-active metal center. An example is represented by system 8 reported in Figure 12.5, constituted by two tetramine coordinating sites, tren [tris(2-aminoethyl)amine)] and cyclam (1,4,8,11-tetraazacyclotetradecane). These binding moieties, separated by a 1,4-xylyl spacer, display different coordinating tendencies toward metals [35]. Consecutive reaction of 8 (Figure 12.5) with one equivalent of Ni(ClO4)2 and one equivalent of Cu(ClO4)2 under the proper conditions leads to the formation of a heterodimetallic complex in which the Ni(II) ion is firmly encircled by the tetra-aza macrocycle, whereas Cu(II) is coordinated by the tripodal tetramine. If one equivalent of tetra-alkylammonium chloride is added to a mM solution of the dimetallic complex, the Cl− ion binds the Cu(II) center to form a five-coordinate complex with an axially compressed trigonal bipyramidal geometry (8a in Figure 12.5).

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Cl N H N

Ni2+

H N H N

N H N

–

Ni3+

H N H N

S

CI

H2N

–

+e– –e–

Cu2+ NH NH2

H2N

Cu2+ NH NH2 N

N

8a

8b

FIgURE 12.5 Redox-driven translocation of the Cl− anion, based on the Ni(III)/Ni(II) redox couple. S represents a solvent molecule.

Spectrophotometric titration experiments enabled to determine the binding constant: log K = 5.66 ± 0.09. Therefore, for a mM solution of 8, addition of one equivalent of anion leads to 95% of the anion bound to the Cu(tren)2+ subunit. The Ni(cyclam)2+ moiety is not competitive for anion binding because Ni(II) prefers the square coordination provided by the cyclam macrocycle. The Ni(II) center undergoes one-electron oxidation at a moderately positive potential [36], and the so-formed d7 cation, in the low-spin state, displays a strong preference toward higher coordination. Thus, upon Ni(II) oxidation, the chloride ion moves from the Cu(II) center to the Ni(III) one, while the free coordination position on the Cu(II) ion is taken by a solvent molecule (8b in Figure 12.5). 3. Helicate Assembly and Disassembly Metal-centered redox processes can also cause a rearrangement of the ligand, as it happens in the case of the assembly and disassembly of metal helicates. For example, compound 9 [37] (Figure 12.6) gives stable complexes with Cu(I) and Cu(II), both in CH3CN solution and in the solid state.

N

N

N

N

9

Cu2+ Cu+ –2e – Cu+

+2e –

+

Cu2+

FIgURE 12.6 Redox-driven assembly–disassembly of a helicate based on Cu(II)/Cu(I) redox couple.

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In particular, Cu(I) gives rise to a dinuclear complex, [Cu2(9)2](CF3SO3)2, in which each Cu(I) ion is bound to an imine and to a pyridine nitrogen atom from each strand, and shows a rather distorted tetrahedral coordination geometry. On the other hand, Cu(II) complexation by ligand 9 leads to a mononuclear complex, [Cu(9)](CF3SO3)2, that reaches tetragonal coordination through chelation by a single molecule of 9 in order to better profit from LFSE terms. Both these complexes were analyzed by x-ray diffraction studies. In solution, assuming that the previously described geometrical features are maintained, the Cu(II)/Cu(I) redox change would result in an assembly–disassembly equilibrium, as pictorially illustrated in Figure 12.6. The occurrence of the redox-driven reversible assembly–disassembly process involving copper complexes of 9 has been verified through cyclic voltammetry experiments at a platinum electrode in CH3CN solution. Both electrochemical processes illustrated in Figure 12.6 appear to be chemically irreversible because of fast assembly–disassembly processes. The analysis of the electrochemical data suggests that the transient species [Cu2(9)2]4+ and [Cu(9)]+ have a lifetime lower than 50 ms. It has been demonstrated that the lifetime of the dicopper(II) double-strand helicate [Cu2(9)2]4+ can be significantly increased by introducing hindering substituents on the framework of 9 [38].

C.

HYDROPHOBIC INTERACTIONS

Cyclodextrins (CDs) [39] and cucurbit[n]urils (CB[n]) [40] represent two classes of host molecules relatively soluble in water that contain a rather rigid and well-defined cavity that can encapsulate guests mainly by hydrophobic interactions. In spite of these general similarities, there are also important differences between them: the cavity of the CDs reminds the shape of a truncated cone, reaching its maximum diameter at the wider opening, while that of the CB[n] has the shape of a barrel and its maximum diameter is located at the molecular equator. Moreover, negative charge density (from the carbonyl oxygens) accumulates on the CB[n] cavity openings, while the hydroxyl groups on the CD portals are believed to be extensively hydrogen bonded. The inclusion complexes of CDs are characterized by association constant up to 103–105 M−1. The formation of highly stable CB[n] inclusion complexes is usually driven by a combination of hydrophobic forces and ion–dipole interactions between strategically located positive charges on the guest and the rims of carbonyl oxygens on the cavity portals [40]. As a result, CB[n] inclusion complexes in aqueous solution may reach association constants as high as 1015 M−1 (equivalent to that of the avidin–biotin host–guest pair [41]). The first work in this area reported the binding interactions of ferrocenecarboxylate anions (FcCOO −) with β-CD in aqueous media buffered at pH 7 (Figure 12.7) [42]. The formation of a stable inclusion complex was demonstrated by electronic absorption spectroscopy. In the presence of β-CD, the half-wave potential (E1/2) for the one-electron oxidation of FcCOO − shifts to more positive values. FcCOO − is more stabilized by β-CD than its mono-oxidized counterpart, Fc+COO −. An additional observation is that the current levels associated with the voltammetric wave decrease in the presence of β-CD. This is easily explained by the larger molecular weight and size of the [β-CD•FcCOO]− complex compared to the free guest. Therefore, inclusion complexation tends to slow down the diffusional flow of guest to the electrode surface. This flow is driven by the low concentration of FcCOO − near the electrode, where it is consumed (oxidized) by the electron transfer process. FcCOO − forms a stable inclusion complex with the host, while the oxidized, zwitterionic form of the guest does not interact strongly with β-CD. The electrochemical oxidation process is best rationalized by an electron transfer process preceded by a coupled chemical reaction (host–guest complexation/decomplexation equilibrium), which is usually referred to as a chemical–electrochemical (CE) mechanism. It is important to realize that the mechanism does not contemplate the direct electrochemical oxidation of the inclusion complex (right process in Figure 12.1). Evidence for this surprising finding was gathered at fast scan rates, at which the

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COO COO

–

–

Fe

Fe

+ β-CD

+e–

–e–

COO

–

Fe +

FIgURE 12.7 Chemical–electrochemical (CE) mechanism for the one-electron oxidation of FcCOO − in the presence of β-CD.

cyclic voltammetric wave for oxidation of FcCOO − is flattened [42], because the complex dissociation mechanism becomes too slow to generate enough free guest to sustain the fast electrochemical oxidation. The lack of electrochemical reactivity of the inclusion complex is linked to the fact that the electrochemical kinetics of the inclusion complex is slower than that of the free guest (see Section II.B) and, because the inclusion complex is short-lived, dissociation provides a bypass mechanism for the electron transfer to take place more quickly. Reinhoudt and coworkers have taken advantage of the multivalent interactions between dendrimers with multiple surface ferrocene residues and SAMs containing CD-binding sites for the development of molecular printboards (see Section III.E). Further investigations of the electrochemical behavior of CD complexes suggest that their labile kinetic nature allows relatively efficient electron transfer reactions to or from the free guest, following the fast dissociation from the host. Therefore, in spite of their accessibility and simplicity of use, these hosts do not form complexes of enough kinetic stability to allow the investigation of the electron transfer reactions involving the encapsulated guests. The formation of a stable inclusion complex between the organic dication N,N′-dimethyl-4,4′bipyridinium (methyl viologen or paraquat, MV2+ [43]) and CB7 was reported for the first time in 2002 [44]. In aqueous media, the stability constant of this complex is ca. 105 M−1, although this value decreases as the ionic strength of the medium increases [45]. The lifetime of the complex was estimated to be 5.3 ms. In aqueous solution, one- or two-electron reduction of MV2+ leads to less charged or neutral species, which are more hydrophobic and have a marked tendency to precipitate on the electrode surface. However, if the discussion is limited to the first one-electron reduction (MV2+ → MV+), one can see that the presence of CB7 decreases the current levels and shifts the half-wave potential to slightly more negative values (Figure 12.8). This finding shows that CB7 differentially stabilizes the dicationic viologen form versus the cation radical, although the small magnitude of the CB7-induced, half-wave potential shift (ca. 30 mV) suggests only a moderate stability drop in the complex after the viologen guest loses one of its positive charges. The decrease in the current levels in the presence of CB7 reflects the larger size and lower diffusivity of the complex compared to the free guest.

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2.0 × 10–5 A

(a)

–0.2 (b)

–0.4

–0.6

–0.8

E (V vs Ag/AgCl)

FIgURE 12.8 Scan rate dependence of the cyclic voltammetric response on a glassy carbon electrode (0.072 cm2) of 1.0 mM MV2+ in 0.1 M NaCl (a) in the absence and (b) in the presence of 1 equiv CB7. Scan rates: 0.1, 0.2, 0.3, 0.4, 0.5, 0.6, 0.8, and 1.0 V s−1 (arrows indicate the increasing scan rates). Potential is referred to Ag/AgCl. (From Ceroni, P., Credi, A., Venturi, M., and Schalley, C.A. eds.: Analytical Methods in Supramolecular Chemistry. pp. 371–457. 2012. Copyright Wiley-VCH Verlag GmbH & Co. KGaA. Reproduced with permission.)

An interesting feature of the [CB7•MV]2+ system is that, in clear contrast to CD-based species, the voltammetric data do not contain any indication that complex dissociation must precede any of the electron transfer processes. Moreover, the electrochemistry of the inclusion complex is as fast—in the timescale accessible in these CV experiments—as that of the free guest. This is clearly illustrated by the voltammograms depicted in Figure 12.8, which shows the comparative results of a scan rate study on the MV2+/MV+ and [CB7•MV]2+/[CB7•MV]+ redox couples. In both cases, the observed anodic and cathodic peak potentials are basically invariant as the scan rate is increased up to 1.0 V/s, revealing the reversible character of both couples under the experimental conditions of the study. Therefore, it was concluded that the inclusion of methyl viologen into CB7 does not affect the electrochemical kinetics of the former in a pronounced way. Although no crystal structure of the [CB7•MV]2+ complex is available, all the experimental and computational data are consistent with a structure in which the two positively charged nitrogens interact with the carbonyl rims on the cavity portals and hydrophobic interactions further stabilizing the midsection of the complex. Obviously the guest is not fully encapsulated, as each of its N-methyl termini protrudes through the host portals.

D.

CHARGE-TRANSFER INTERACTIONS

The interaction between electron-poor and electron-rich molecules, particularly those containing π-systems, is a common driving force for the formation of host–guest systems. An interesting class of hosts is constituted by molecular clips and tweezers [46–48], which have a well-organized, yet relatively flexible shape that enables interactions with guests of different sizes. The recognition process can be based on different kinds of interaction, but an example based mainly on electron donor– acceptor interactions is reported. Molecular clip 10 is constituted by extended aromatic sidewalls

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Electrochemically Driven Supramolecular Devices

(namely, benzo[k]fluoranthene units [BkF]) connected by a dimethylene-bridged spacer containing a benzene ring that carries two acetoxy groups. The BkF moieties are strongly fluorescent [49] polycyclic aromatic hydrocarbons with good electron donor ability [50]. The resulting clip is soluble in organic solvents and presents a highly negative electrostatic potential on the inner van der Waals surface of the cavity that forms a recognition site for electron-deficient guest molecules. A highly stable host–guest complex of clip 10 with 9-dicyanomethylene-2,4,7-trinitrofluorene (TNF) has been reported [51]: K ≈ 105 M−1, in chloroform solution, estimated by spectrophotometric and spectrofluorimetric titrations. R NC

CN

R O2N NO2

10 O R=

NO2

TNF

O

As to the electrochemical properties, clip 10 presents two one-electron transfer processes associated with chemical reactions, corresponding to the first oxidation of each benzo[k]fluoranthene sidewall that occurs at distinct potential values because of electrostatic and/or electronic interactions between the two moieties. The oxidation of BkF is known to be associated with a dimerization reaction [52]. TNF shows three reversible one-electron reduction processes: the first process occurs at −0.02 V (vs the saturated calomel electrode [SCE]), demonstrating the strong electron acceptor character of this compound [53]. An equimolar solution of 10 and TNF (1.3 mM) in CH 2Cl 2/TBAPF6 shows cyclic voltammograms in which the first oxidation process is positively shifted and the first reduction process is negatively shifted, compared to the CVs of the two separated components. Under the experimental conditions used, the fraction of [10•TNF] formed is larger than 90%. The fact that the benzo[k]fluoranthene moiety is more difficult to oxidize and TNF is more difficult to reduce when they are associated with one another demonstrates the charge transfer character of this complex. Moreover, lower current intensities are observed for [10•TNF] because of the smaller diffusion coefficient of the complex compared to the separated components. On the other hand, successive reduction and oxidation processes occur at the same potentials of the isolated components, because one-electron oxidation or one-electron reduction causes the disruption of the complex. By successive addition (from 0.5 to 3 equivalents) of clip 10 to a 1.2 mM solution of TNF, a progressive shift of the first reduction process is observed (Figure  12.9): the maximum shift of the half-wave potential compared to the free TNF is 95 mV, recorded upon addition of two equivalents of clip per guest molecule. The observed electrochemical behavior can be rationalized on the basis of the square scheme reported in Figure 12.1, in which the association constant K of 10 and TNF is much larger than that of 10 and TNF− (K′), giving a K/K′ ratio of the order of 40. Upon variation of the scan rate in the range 0.05–20 V/s, no change in the shape of the cyclic voltammogram is observed, indicating that both the forward and backward scans of the first reduction process are due to the redox couple [10•TNF]/[10•TNF]−. The heterogeneous electron transfer rates are, in any case, higher than the association and dissociation rate constants of the complex between 10 and TNF or 10 and TNF−. A special class of host–guest systems is represented by pseudorotaxanes [54], in which the guest has a threadlike shape and the host a ring structure, so that the extremities of the thread are directed away from the center of the host. At least one of the extremities of the thread does not have

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Fc 10

i (μA)

0 –0.3

0.0

0.3

0.6 E (V vs SCE)

–10

–20

FIgURE 12.9 Cyclic voltammograms of TNF (1.2 mM) upon addition of clip 10 from 0 (thick solid line) up to 2.0 (dashed line) equivalents in CH2Cl2/TBAPF6 solution. Working electrode: glassy carbon. Scan rate: 0.05 V/s. The reversible wave of ferrocene (Fc) used as an internal standard is also reported.

a bulky stopper group. Hence, the constituents of the assembly, like any complex, are at liberty to dissociate into separate molecular species (i.e., in contrast to rotaxanes, see Section IV.B, there is no attendant mechanical bond to maintain the integrity of the system). Pseudorotaxanes can be based on different types of intermolecular forces; in this section, only systems characterized by the presence of electron donor–acceptor interactions are described. One of the most extensively studied receptors in the past two decades has been the cyclophane 114+, which is a very efficient host for a wide variety of π-electron-donating guests [55].

+N

N+

+N

N+

114+

Two bielectronic reduction processes are observed for this tetracationic cyclophane: the first corresponding to the uptake of the first electron by each of the equivalent bipyridinium units and the second one to the subsequent reduction of the radical cations to neutral units [56]. When an electron donor unit is located inside the cavity of the cyclophane, the potential associated with the first reduction process is shifted to more negative values, as a consequence of the charge-transfer interactions that stabilize the complex [16]. The second reduction process at more negative potentials of this cyclophane is very important, because it can be used to monitor the occurrence of decomplexation induced by the first two-electron reduction [16], as in the case of the previously described [10•TNF] complex. For example, in the presence of excess of a threadlike compound composed of a polyether chain, which bears a 1,4-dioxybenzene unit in the middle, the potential value for the first bielectronic reduction of 114+ is shifted cathodically, whereas the second reduction process is almost unaffected [56]. This observation is consistent with the formation of a pseudorotaxane between the cyclophane and the thread, and dethreading of the pseudorotaxane upon two-electron reduction of the 114+ host, so that the second two-electron reduction process reflects that of the

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Electrochemically Driven Supramolecular Devices

free host. The occurrence of the dethreading reaction is not surprising, because reduction of the electron-acceptor component weakens the charge-transfer interaction that helps to hold together the components of the supramolecular architecture. Because all these processes are reversible, oxidation back to the tetracationic form affords the original pseudorotaxane. It should, in principle, also be possible to obtain useful information about the occurrence of dethreading–rethreading processes from the electrochemical behavior of the guest; the poor reversibility of the oxidation process associated with a 1,4-dioxybenzene unit, however, prevents the use of this type of control. More interesting are pseudorotaxanes wherein both the cyclophane and thread components are characterized by chemically reversible redox processes; one example is the complex of TTF with 114+ [57] and related pseudorotaxanes [57a,58]. This improvement in design not only enables to monitor the formation of the supramolecular species by studying both the reduction of the electron acceptor component and the oxidation of the electron donor species but also provides a dual mode (reductive and oxidative) of control on the dethreading–rethreading processes. The molecular thread 12, obtained by attaching two polyether chains to a TTF unit (Figure 12.10), forms a very stable pseudorotaxane with 114+. O

O OH

+

S

O O

+

+

N

O

S

N O

O

+

S

N

O

S

S

+

O

N+

S

Dethreading O

S

O N +

+

N

+

S

O HO

HO +e–

+

+

N S

S

O O

S

N O

O

N

O

S

O N +

N

O

(Re)threading O

S S

S

+

S

O

N+

N

O

HO

HO

114+

O

–2e– O

O N

O O

S

O

12

+

S

O

N+

+

+2e–

O

OH

O

+

+

[11 12]4+

OH

O

O

O

O

O

N+

+N

–e–

OH

OH

O

S

N O

O

S

O N

N

O

O N+

HO

+

N

Dethreading

OH

O O +

S S

S S

N

+

O

N O HO

O

O

FIgURE 12.10 Redox-driven (re)threading–dethreading of a pseudorotaxane stabilized by electron donor– acceptor interactions.

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Reversible dethreading–rethreading cycles of the pseudorotaxane [11•12]4+ can be performed either by oxidation and successive reduction of the electron-donating thread or by reduction and successive oxidation of the electron-accepting cyclophane. Such processes are accompanied by pronounced visible spectral differences that can be followed easily by the naked eye. Interestingly, while monoelectronic oxidation of TTF to its radical cation cancels the strong electron donor character, its two-electron oxidation renders the formed TTF2+ dication a moderate electron acceptor [59]. This property was exploited to devise a redox-controlled three-pole switch [60] based on a supramolecular system composed of TTF and two hosts, namely, the π-electronaccepting cyclophane 114+ and the π-electron-donating crown ether 13 (Figure 12.11) [61]. In its role of an electron donor, TTF forms, in acetonitrile solution, a 1:1 inclusion complex with 114+ that can be dissociated–reassociated reversibly by cyclic oxidation–reduction of TTF, while TTF2+ acts as a π-electron acceptor, giving a 1:1 inclusion complex with 13. In contrast, TTF+ is not bound by either of the two hosts. When the electrochemical potential applied to the solution, where TTF is complexed by 114+, [11•TTF]4+, becomes more positive than ~ +0.4 V (relative to the SCE), TTF is oxidized to the monocation form and the complex disassembles to give three essentially noninteracting species. Further one-electron oxidation to TTF2+ at potentials more positive than ~ +0.7 V (relative to the SCE) leads to the insertion of the dication into the cavity of 13, [13•TTF]2+. Because both oxidized forms of TTF are stable, the initial state can be restored by subsequent reduction. This system can therefore be switched reversibly between three distinct states by exercising electrochemical control on the guest behavior of TTF (Figure 12.11). The fact that the three states have different colors, coupled with the ease of their electrochemical interconversion, renders this supramolecular system suitable for electrochromic applications; the system could, moreover, form the basis for the construction of molecular devices in which energy or electron transfer processes between selected components can be controlled [61]. It has been shown more recently [62] that the tendency for bipyridinium radical cations to dimerize can be exploited to assemble strong complexes between the dicationic cyclophane host 112+ (obtained upon two-electron reduction of 114+) and bipyridinium radical cationic guests (obtained upon one-electron reduction of the corresponding dications). Such a recognition motif forms the basis of redox-switchable supramolecular species and mechanically interlocked molecules [62–64].

N+

+N

+

N+

N

114+

+N

+N

N+

O O S

+

N+

S

S

S

S

+N

+e–

S

–e–

S

+ O O

S

TTF+

S

+

S O

N+

+e–

+

O

N+

+N

–e–

O

O

O

O O

O

O O

O

S+S O

O O

O

O

O O

O O

O 13

O

O O

O

O

FIgURE 12.11 A supramolecular switch in which selection of either cyclophane 114+ or crown ether 13 hosts is achieved by changing the oxidation state of the TTF guest.

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Electrochemically Driven Supramolecular Devices

E.

HOST-gUEST SYSTEMS WORKING ON SURFACES

The three-pole supramolecular switch described in Section III.D (Figure 12.11) was also constructed and operated at the solid–liquid interface of an electrode [65]. A SAM of compound 14 was obtained on a gold electrode. Voltammetric experiments showed that, depending on the oxidation state of the immobilized guest unit, the TTF moiety can form a complex either with cyclophane 114+ (when it is in the neutral TTF form) or with crown ether 13 (when it is in the dicationic TTF2+ form). Hence, the multistage redox properties of SAMs of 14 can be exploited to produce surfaces with electrochemically controllable binding properties. Unfortunately, the reversibility associated with the sequential oxidation of the TTF moiety of the SAM in the presence of both 114+ and 13 macrocycles was poor, which will significantly hamper the fabrication of usable devices from this system. O S

S

S

S

O S

S

14

Water-soluble supramolecular assemblies of redox-active ferrocenyl-decorated dendrimers such as 15 (Figure 12.12) and β-CD were adsorbed at molecular printboards [27] composed of SAMs of β-CD [66]. The dendrimers form a stable monolayer at the β-CD SAM owing to multivalent host–guest interactions. The immobilization of the dendrimers can be electrochemically controlled

Fe

Fe

O HN Fe

HO

O

O O

H N

N

Fe

H N

N

O

HN O

7

O N

N

O Fe

OH

NH

O N

N H

N

N H

Fe HO

HN

OH

NH O

15

O

S

O

O

Fe

Fe

(a)

HO

7

β-CD

Oxidation + +

+

(b)

FIgURE 12.12 (See color insert.) (a) Structures of the ferrocene-terminated dendrimer 15, β-cyclodextrin, and the β-cyclodextrin derivative used for the preparation of the SAMs. (b) Schematic representation of the self-assembly of the dendrimer-β-CD complex, its adsorption on the β-CD modified surface, and the electrochemically induced desorption.

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because, while ferrocene forms stable inclusion complexes with β-CD, the ferrocinium cation does not (Figure 12.12). The redox-dependent adsorption process was studied by voltammetric techniques and electrochemical impedance spectroscopy. Detailed information was obtained on the adsorption and desorption kinetics and mechanism, diffusion coefficients, and conformation of the guest dendrimers adsorbed at the host surface [66]. More recently, the electrochemical modulation of host–guest interactions has been used to control the aggregation–disaggregation of nanoparticles [67].

IV. MOLECULAR MACHINES A. BASIC CONCEPTS The design, synthesis, and operation of molecular-scale systems that exhibit controllable motion of their component parts is a topic of great interest in nanoscience and a fascinating challenge of nanotechnology. The development of this kind of species constitutes the premise to the construction of molecular machines and motors, which in a not too distant future could find applications in fields such as materials science, information technology, energy conversion, diagnostics, and medicine [8,68]. Like the macroscopic counterparts, molecular machines can be classified according to characteristics such as the type of energy supply to make them work, the kind of motion (translation, rotation, oscillation, etc.) performed by their components, the way of monitoring their operation (the rearrangements of the component parts should cause readable changes in some chemical or physical property of the system), the possibility to repeat the operation in cycles, the timescale needed to complete a cycle, and the function that can be ultimately carried out. It should be recalled, however, that nanoscale machines cannot be simply considered as shrunk versions of the macroscopic counterparts because several intrinsic properties of molecular-level entities are quite different from those of macroscopic objects. A thorough discussion on the basic concepts of molecular machines can be found elsewhere [8,68]. Molecular machines operate through chemical reactions and need to be fed by an energy source. The obvious way to provide a chemical system with energy is through an exoergonic chemical reaction. If a molecular machine has to work by inputs of chemical energy, addition of fresh reactants (fuel) at any step of its working cycle is needed, with the concomitant formation of waste products. Research in the field has shown that artificial molecular machines can also conveniently be powered by electrical energy (through electrochemically induced redox reactions) and by light (through photoinduced reactions such as photoisomerization and photoinduced electron transfer processes). In the past 25 years, the development of supramolecular chemistry has enabled the construction of an interesting variety of artificial molecular machines. Most of them are based on the topologically intriguing chemical systems called rotaxanes and catenanes. Simple examples of molecular machines are also pseudorotaxanes (see Section III.D) because of the possibility of controlling the component threading/dethreading motion.

B.

SYSTEMS BASED ON ROTAXANES

Rotaxanes [69], the name of which derives from the Latin words rota and axis for wheel and axle, are minimally composed (Figure 12.13) of a macrocyclic compound (the ring) threaded by a dumbbell-shaped molecule terminated by bulky groups (stoppers) that prevent disassembly. Important features of these systems derive from noncovalent interactions between components that contain complementary recognition sites. The switch off and again on of such interactions is at the basis of the machine-like behavior of rotaxanes as schematized in Figure 12.13a and b.

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Electrochemically Driven Supramolecular Devices

451

(a)

(b)

FIgURE 12.13 Schematic representation of a rotaxane structure and the simple movements that can occur upon appropriate energy stimulation: (a) ring shuttling and (b) ring pirouetting.

1. Molecular Shuttles When a rotaxane contains two different recognition sites in its dumbbell component, it can behave as a controllable molecular shuttle (Figure 12.13a), and, if appropriately designed by incorporating suitable redox units, it can perform its machine-like operation by exploiting electrochemical energy inputs. Of course, in such cases, electrons/holes, besides supplying the energy needed to make the machine work, can also be useful to read the state of the system by means of the various electrochemical techniques. Three paradigmatic examples of rotaxanes as electrochemically driven shuttles are reported, which differ for the nature of the interactions stabilizing their structure: electron donor–acceptor interactions for the first system, hydrogen-bonding and π-stacking interactions for the second one, and strong metal–ligand bonding for the third one. Rotaxane 16 6+, specifically designed [70,71] to achieve photoinduced ring shuttling in solution [72], behaves also as an electrochemically driven molecular shuttle (Figure 12.14). Because of the presence of several redox-active units, the cyclic voltammogram of this rotaxane shows a complex redox pattern. However, the comparison with the electrochemical behavior of its molecular components and suitable model compounds (Figure 12.14) enables to obtain useful information not only on its conformational features, but also, and most importantly, on its machine-like operation. In dumbbell-shaped component 18 6+, all the redox processes of the incorporated units are present at almost the same potentials as in the separated units (Figure 12.15a); this finding shows that there are no substantial intercomponent electronic interactions. On going from the dumbbell component to rotaxane 16 6+, some processes are affected while others are not (Figure 12.15a). All the processes related to the Ru-based unit, namely, the metal-localized oxidation and the ligand-localized reductions, do not show any appreciable changes. A different behavior is, however, observed for the A12+ unit, the first reduction of which is displaced noticeably toward more negative potential values. This result indicates that it is surrounded by the electron donor macrocycle 17 (Figure 12.15b). Accordingly, the oxidation of the two dioxybenzene (DOB) units of the macrocycle is displaced toward more positive potential values and occurs simultaneously (Figure 12.15b), as already observed for other rotaxanes containing ring 17 [73]. The fact that the ring encircles the A12+ station, as confirmed by the NMR spectrum, is an expected result on the basis of the reduction potentials of A12+ and A22+ in component 186+ (or of separated model compounds). The second process of 16 6+, which corresponds to the first reduction of the A22+ station, is also displaced toward more negative potential values (Figure 12.15b), demonstrating that, at this stage, the A22+ unit is encircled by macrocycle 17. A further proof of the ring displacement is given by the fact that the second reduction of the A12+ station occurs practically at the same potential of the dumbbell (Figure 12.15b). This behavior confirms that, when the better station (A12+) of the two has been deactivated upon reduction, the ring moves to the alternative A22+ station. Under these conditions, from the values of the first reduction potential of A22+ and the second reduction potential of A12+ in the dumbbell

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452

Organic Electrochemistry P2+ N

S N

+

O

O O

+

N

N Ru2+ N N

A22+

O A1

+

O

N

T

O

+

N

N

R

O

2+

N

O

O

O

O

O

O

O

O

O

166+ O O

N

N

+

N

N

N

N

O

+

+

N

N

N

O

17

O

N Ru2+ N

O

O

O

+

N

186+

+

H3C N

OH

+

N CH3

222+

N Ru2+ N N

N

+

192+

+

N

N

+

N

+

N

O

O

204+ +

N+

N 212+

FIgURE 12.14 Structure formulas of rotaxane 16 6+, its ring and dumbbell-shaped components 17 and 18 6+, respectively, and model compounds 192+, 204+, 212+, and 222+ of the units present in the dumbbell.

component, it can be estimated that the translational isomer with the ring surrounding A22+ is much more populated than that in which the ring encircles A1+. When also the A22+ station has been monoreduced, the position of the ring is no longer controlled by strong CT interactions; from the electrochemical results, it seems that it resides close to A2+. The reversibility of the electrochemical processes involving the two stations shows that, after a twoelectron reduction of rotaxane 16 6+, one-electron oxidation relocated the ring on the A22+ station and a successive one-electron oxidation entices it back again onto the A12+ station. The chemical and electrochemical reversibility of these processes also indicates that the rates of the electrochemically induced ring movements are fast. The second example is represented by rotaxane 23 (Figure 12.16), which comprises a macrocycle containing aromatic hydrogen–bonding amide sites, and an axle containing a hydrogen–bonding station (a glycylglycine unit) and a fullerene stopper separated by a triethylene glycol spacer [74]. As expected, in solvents such as CH2Cl2, CHCl3, and THF, the macrocycle stays preferentially on the peptidic station, far away from the fullerene. Instead, in solvents such as DMSO and DMF that weaken the hydrogen bonds, the interactions between the macrocycle and the fullerene are promoted, inducing a large positional change of the macrocycle. In the excited state absorption measurements, nearly solvent-independent behavior was observed for the thread, while measurements carried out on the rotaxane showed that the fluorescence of the fullerene is quenched by

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Electrochemically Driven Supramolecular Devices

17 ×0.5 166+

186+

192+

204+

×0.5 20 μA

212+ ×0.5 222+

+1.5

+1.0

+0.5

(a)

0 –0.5 E (V vs SCE)

–1.0

–1.5

–2.0

186+

17

166+ +1.4 +1.2 E (V vs SCE)

–0.4

–0.6 –0.8 E (V vs SCE)

–1.0

(b)

FIgURE 12.15 (a) Cyclic voltammetric patterns for rotaxane 16 6+, its ring 17, and dumbbell-shaped component 18 6+ and model compounds 192+, 204+, 212+, and 222+ in CH3CN solution. (b) Potential shifts caused by the donor–acceptor interaction between ring 17 and dumbbell-shaped component 18 6+ when they are assembled in rotaxane 16 6+. Circles, squares, and triangles represent processes centered on dioxybenzene, A12+ and A22+ units, respectively. (From Ceroni, P., Credi, A., Venturi, M., and Schalley, C.A. eds.: Analytical Methods in Supramolecular Chemistry. pp. 371–457. 2012. Copyright Wiley-VCH Verlag GmbH & Co. KGaA. Reproduced with permission.)

the proximity of the macrocycle by 28% and 44%, respectively, in DMF and DMSO. The residual fullerene fluorescence in DMSO is 51% of that in CH2Cl2. These experiments provided not only a way to identify the interactions taking place between the macrocycle and the fullerene but also a much simpler way to monitor shuttling than with transient absorption measurements. An important feature of rotaxane 23 is that the position and the translocation of the macrocycle can also be monitored and achieved by the reduction of the fullerene to its trianion, which is both affected and observed by cyclic voltammetry. In DMSO, the proximity of the macrocycle to the

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454

Organic Electrochemistry

O

O

23 N

N HH O N

O

O

H N

N H

N

O H H N

O CHCl3

O

DMSO

+3e–

–3e–

O O HN

O N

O

O

H N

N H

O

O

HN O

FIgURE 12.16 Conformational rearrangements in rotaxane 23 induced by changing the solvent polarity and by electrochemical reduction of the fullerene stopper.

fullerene stabilized substantially the electrogenerated trianion (ΔE1/2 = 46 mV) through π–π interactions. Surprisingly, in THF where the macrocycle is preferentially positioned on the peptide station, a similar behavior was observed (ΔE1/2 = 40 mV). Although contradictory, this result can be easily rationalized in terms of an electrochemically induced shuttling of the macrocycle from the peptide station to a position near the fullerene where it is stabilized by the negative charge present on this stopper. The possibility of fine-tuning electron transfer processes through molecular shuttling was finally shown introducing ferrocene electron donors on the macrocycle (rotaxane 24) [75]. Fe O

O

O N N

O

O

O N

HH O H N N

H N

O H H N

O

O O

24

O Fe

Steady-state and transient absorption photophysical measurements revealed through-space photoinduced electron transfer between the fullerene stopper and the ferrocenes on the macrocycle.

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Electrochemically Driven Supramolecular Devices

In CH2Cl2, the radical ion pair state lifetime of 26.2 ns was measured. This value is consistent with the larger relative separation of the electroactive units that give longer lifetimes. Indeed, the use of a polar solvent (hexafluoroisopropanol) shortens the lifetime to 13.0 ns, a consequence imposed by weakening the hydrogen bonds that decrease the relative spatial separation between the donor and the acceptor, while increasing the shuttling rate. The third example of electrochemically driven shuttles is rotaxane [25•Cu]+ (Figure 12.17) that is stabilized by metal–ligand bonding. It contains a phenanthroline and a terpyridine unit in its dumbbell-shaped component; it also incorporates a Cu(I) center coordinated tetrahedrally by the phenanthroline ligand of the dumbbell together with the phenanthroline ligand of the macrocycle [76–78]. Oxidation of the tetracoordinated Cu(I) center of [25•Cu]+ to a tetracoordinated Cu(II) ion occurs on electrolysis (+1.0 V relative to the SCE) of a CH3CN solution of the rotaxane. In response to the preference of Cu(II) for a pentacoordination geometry, the macrocycle shuttles away from the bidentate phenanthroline ligand of the dumbbell and encircles the terdentate terpyridine ligand instead. O

+

O

O

O

O

O N N

N Cu N N

N

N

O

–

O

[25 Cu]+

O

– e–

–

Cu

–

Cu

+e–

–

Cu [25 Cu]2+

FIgURE 12.17 Shuttling of the macrocyclic component of [25•Cu]+ along its dumbbell-shaped component controlled electrochemically by oxidizing–reducing the metal center.

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Organic Electrochemistry

Consistently, the cyclic voltammogram shows the disappearance of the reversible wave (+0.68 V) associated with the tetracoordinated Cu(II)/Cu(I) redox couple and the concomitant appearance of a reversible wave (−0.03 V) corresponding to the pentacoordinated Cu(II)/Cu(I) redox couple. A second electrolysis at −0.03 V of the acetonitrile solution of the rotaxane reduces the pentacoordinated Cu(II) center back to a pentacoordinated Cu(I) ion. In response to the preference of Cu(I) for a tetracoordination geometry, the macrocycle moves away from the terdentate terpyridine ligand and encircles the bidentate phenanthroline ligand. The cyclic voltammogram recorded after the second electrolysis shows the original redox wave (+0.68 V) corresponding to the tetracoordinated Cu(II)/Cu(I) redox couple. This system has been improved by replacing the highly shielding and hindering phenanthroline moiety contained in the ring with a nonhindering bi-isoquinoline unit [79]. In the new rotaxane, the electrochemically driven shuttling of the ring is, indeed, at least four orders of magnitude faster than in the previous phenanthroline-based system. 2. Ring Pirouetting Motion In suitably designed rotaxanes, the pirouetting-type movements of the ring around the axle (Figure 12.13b) can be electrochemically driven. Rotaxane [26•Cu]+ has a structure (Figure 12.18a) in which Cu(I) is coordinated tetrahedrically by the phenanthroline present in the axle and the phenanthroline contained in the ring [80,81]. Electrochemical oxidation of the Cu(I) center leads to a transient tetracoordinated Cu(II) species that, by the pirouetting of the ring around the axle, rearranges in tens of seconds to a structure in which the Cu(II) center reaches its most stable environment, being pentacoordinated by the phenanthroline of the axle and the terpyridine of the ring ([26•Cu]2+ in Figure 12.18b). On electrochemical reduction of Cu(II), a transient pentacoordinated Cu(I) species is obtained, which rearranges in the millisecond timescale by means of a second pirouetting of the ring to the most stable structure with Cu(I) tetrahedrically coordinated ([26•Cu]+ in Figure 12.18b). The obtained results underline that the rearrangement rates from the less to the most stable geometries are drastically different for the two oxidation states of the metal. In order to increase the rate of the ring pirouetting, the new rotaxane [27•Cu]+ (Figure 12.18a) was prepared [82] in which the metal center is more accessible because of the presence of a less hindered ligand compared to the previous related system. Ligand exchange within the coordination sphere of the metal is thus facilitated as much as possible. The molecular axle contains indeed a thin 2,2′-bipyridine motif, which is less bulky than a 1,10-phenanthroline fragment and it is expected to spin more readily within the cavity of the ring as a consequence of Cu oxidation/reduction. Compared to [26•Cu]+, in this rotaxane, both the oxidation-induced and the reduction-induced ring pirouetting movements are more than three orders of magnitude faster. Such an example shows that subtle structural factors can have a very significant influence on the general behavior (rate of the movement, in particular) of Cu(II/I)-based molecular machines. Further improvement in the pirouetting rate has then been obtained by keeping the two stoppers of [27•Cu]+ very remote from the copper center [83]. The results obtained show that in this improved system, the ring moves very fast around the axle (milliseconds) even at low temperature. It should be noted, however, that, as a consequence of the oxidation–reduction cycle, the ring of these rotaxanes does not necessarily perform a 360° rotation, but can only oscillate between the two positions on the threaded axle. To make real rotary motors, it would be necessary to introduce directionality in the system by using, for example, a ring containing three different coordination sites and a suitably designed axle.

C.

SYSTEMS BASED ON CATENANES

Catenanes [69], whose name derives from the Latin word catena for chain, are made of (at least) two interlocked macrocycles or rings (Figure 12.19a). As for rotaxanes, the important feature of catenanes derives from the presence of weak noncovalent interactions between the components that

© 2016 by Taylor & Francis Group, LLC

457

Electrochemically Driven Supramolecular Devices +

+

O O

N

–

N N

Cu N

N

O O

O N

N N

N N O

O

N

–

O

Cu

N

N N

O O

O O

[26 Cu]+

[27 Cu]+ (a)

–

–

–e– Cu

Cu +e–

(b)

FIgURE 12.18 (a) Structure of rotaxanes [26•Cu]+ and [27•Cu]+ and (b) schematic representation of the ring pirouetting induced by oxidation–reduction of the metal center.

contain complementary recognition sites. Because of these interactions, catenanes, like rotaxanes, are appealing systems for the construction of molecular machines. The large-amplitude motion that can be achieved with catenanes is the circumrotation of one ring with respect to the other, and when suitably designed, they can be seen as simple prototypes of molecular rotors. As already pointed out in the case of rotaxanes, also in catenanes, the dynamic processes of one ring with respect to the other can be controlled if at least two different recognition sites are incorporated in the structure (Figure 12.19b). In particular, if redox units are incorporated in the catenane structure, there is the possibility of controlling these processes upon electrochemical stimulation. An example is offered by catenane 284+ (Figure 12.20a) that incorporates the previously seen macrocycle 17 and a tetracationic cyclophane comprising one bipyridinium and one trans-1,2-bis (4-pyridinium)ethylene unit [84,85].

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Organic Electrochemistry

(a)

+S1

+S2

State 0

(b)

State 1

FIgURE 12.19 Schematic representation of (a) a catenane structure and (b) the two conformational isomers associated with a catenane incorporating two different recognition sites within one of its two macrocyclic components; the two isomers can be interchanged by appropriate stimuli (S1 and S2).

In the major isomer, the bipyridinium unit is located inside the cavity of the macrocyclic polyether and the trans-bis(pyridinium)ethylene unit is positioned alongside, as confirmed by the electrochemical analysis. The cyclic voltammogram of the catenane shows four monoelectronic processes that, by a comparison with the data obtained for the free cyclophane, can be attributed as follows: the first and fourth processes to the first and second reductions of the bipyridinium unit, and the second and third ones to the first and second reductions of the trans-bis(pyridinium)ethylene unit. The comparison with the tetracationic cyclophane also evidences that all these reductions are shifted toward more negative potential values (Figure 12.20b). O O O O O +N

O O O O O +N

N+

N+

+e– N+

+N

28

+N

N+

O O O O O

O O O O O

O O O O O

O O O O O

4+

+N

N+

+N

N+

–e–

+N

N+ O O O O O

N

N+ O O O O O

(a)

FIgURE 12.20 (a) The circumrotation of the tetracationic cyclophane component of catenane 284+ can be controlled reversibly by reducing–oxidizing electrochemically its bipyridinium unit. (Continued)

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459

Electrochemically Driven Supramolecular Devices

Tetracationic ring

Catenane 284+ –0.3 (b)

–0.7 –0.5 E (V vs SCE)

–0.9

FIgURE 12.20 (Continued) (b) correlation between the half-wave reduction potentials of catenane 284+ and of its tetracationic ring component (circles and squares correspond to the reduction of bipyridinium and trans-bis(pyridinium)ethylene units, respectively). (From Ceroni, P., Credi, A., and Venturi, M. eds.: Electrochemistry of Functional Supramolecular Systems. p. 415. 2010. Copyright Wiley-VCH Verlag GmbH & Co. KGaA. Reproduced with permission.)

The discussion can be restricted to the first and second reduction processes that are of particular interest in this context. The shift of the bipyridinium-based process is in agreement with the catenane conformation in which the bipyridinium unit is located inside the cavity of the macrocyclic polyether (Figure 12.20a); because of the CT interactions established with both the electron donor units of the macrocycle, its reduction is more difficult than in the free tetracationic cyclophane. The shift of the trans-1,2-bis(4-pyridinium)ethylene-based reduction indicates that, once the bipyridinium unit is reduced, the CT interactions that stabilize the initial conformation are destroyed and, thereby, the tetracationic cyclophane circumrotates through the cavity of the macrocyclic polyether moving the trans-bis(pyridinium)ethylene unit inside, as shown by comparison of its reduction potential with that of a catenane model compound [84]. The original equilibrium between the two conformations associated with catenane 284+ is restored upon oxidation of both units back to their dicationic states. It is interesting to notice that for a machine-like performance, the presence of a desymmetrized ring is a necessary but not a sufficient requirement. This statement is clearly demonstrated by the behavior of catenane 294+ made by the same tetracationic cyclophane of 284+ and a macrocycle containing two dioxynaphthalene (DON) units [84,85]. As in the case of 284+ the major isomer is the one in which the bipyridinium unit is located inside the macrocycle. However, in contrast with the behavior of 284+, for which the first reduction process concerns the inside bipyridinium unit, the first reduction of 294+ involves the alongside trans-bis(pyridinium)ethylene unit (Figure 12.21).

+N

O O O O O N+ +N

+N

N+

N+ N+

+N

–0.3

–0.5 –0.7 E (V vs SCE)

–0.9

O O O O O 294+

FIgURE 12.21 Correlation between the half-wave reduction potentials of catenane 294+ and of its tetracationic ring component (circles and squares correspond to the reduction of bipyridinium and trans-bis(pyridinium) ethylene units, respectively). (From Ceroni, P., Credi, A., and Venturi, M. eds.: Electrochemistry of Functional Supramolecular Systems. p. 417. 2010. Copyright Wiley-VCH Verlag GmbH & Co. KGaA. Reproduced with permission.)

© 2016 by Taylor & Francis Group, LLC

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Organic Electrochemistry

This  process has been undoubtedly attributed to this unit by performing spectroelectrochemical experiments and comparing the spectrum of the monoreduced catenane with that of a model compound containing only the trans-bis(pyridinium)ethylene unit. The different behavior of the two catenanes, as far as the first reduction process is concerned, can be explained on the basis of the different strength of the CT interactions: in 294+ the bipyridinium unit is sandwiched between two DON moieties that are stronger electron donors than the DOB moieties of the macrocyclic component of catenane 284+. Because of these stronger interactions, the reduction of such a unit becomes so difficult that it occurs at a potential more negative than that of the trans-bis(pyridinium)ethylene unit (Figure 12.21). As a consequence of the fact that in 294+ the first reduction concerns the alongside unit, the CT interactions responsible for the initial conformation are practically unaffected and no mechanical movement occurs in the monoreduced catenane. Catenanes 304+ and 314+ (Figure 12.22) are other examples of systems in which the conformational motion can be controlled electrochemically [58b,86]. They are made of the already seen symmetric tetracationic cyclophane 94+ and a desymmetrized ring comprising two different electron donor units, namely, a TTF and a DOB (304+) or a DON (314+) units. Because the TTF moiety is a better electron donor than the dioxyarene units, as witnessed by the potential values for their oxidation, the thermodynamically stable conformation of these catenanes is that in which the symmetric ring encircles the TTF unit of the desymmetrized one (Figure 12.22a, State 0). The cyclic voltammogram of the free macrocycle shows a reversible process attributed to the monoelectronic oxidation of the TTF unit. In the catenanes, such a process occurs at more positive O O O O O +

+

N

N S

S

S

S

–e–

+N

N+ O

(a)

O

O

O

(b)

O

State 0

+e–

(c)

(d)

State 1

=

304+

314+

FIgURE 12.22 Redox-controlled ring rotation in solution for catenanes 304+ and 314+, which contain the symmetric electron acceptor cyclophane 94+ and a desymmetrized electron donor ring.

© 2016 by Taylor & Francis Group, LLC

Electrochemically Driven Supramolecular Devices

461

potentials, in agreement with the fact that the TTF unit is located inside the cavity of the tetracationic cyclophane and, therefore, engaged in strong CT interactions. It is also interesting to notice that in the catenane, this oxidation process is characterized by a large separation between the anodic and cathodic peaks, which varies as the scan rate is changed. Upon increasing the scan rate, the anodic peak moves to more positive potentials, while the cathodic one shifts to less positive values. These observations indicate that the oxidation–reduction of the TTF unit is accompanied by the circumrotation of the desymmetrized ring through the cavity of the tetracationic cyclophane and that this change is occurring on the timescale of the electrochemical experiment. Indeed, after oxidation, the newly formed monocationic TTF unit (Figure 12.22b) loses its electron donor power; as a consequence, it is expelled from the cavity of the tetracationic cyclophane and is replaced by the neutral dioxyarene unit (Figure 12.22c, State 1). After reduction, the original conformation is restored as the neutral TTF unit replaces the dioxyarene unit inside the cavity of the tetracationic cyclophane. Ring rotation in these catenanes can also be obtained chemically. The tendency of o-chloroanil to stack against TTF has been indeed exploited [58b,86] to lock this unit alongside the cavity of the tetracationic cyclophane. On addition of a mixture of Na 2S2O5 and NH4PF6 in H2O, the adduct formed between the TTF unit and o-chloroanil is destroyed, and the original conformation with TTF inside the cavity of the tetracationic cyclophane is then restored. Catenane 314+ was also incorporated in a solid-state device that could be used for random access memory (RAM) storage. Additionally, this compound could be employed for the construction of electrochromic systems, because its various redox states are characterized by different colors [58b,86,87]. The six electroactive catenanes 322+–372+ (Figure 12.23) [88] are formed by (1) a desymmetrized electron acceptor ring containing two different units, namely, a 4,4′-bipyridinium dication (BPY2+) and a neutral naphthalene di-imide (NDI) or pyromellitic di-imide (PMI) moiety, and (2) an electron donor ring that can be symmetric for the presence of two identical DOB or DON units, or desymmetrized by the presence of two different donors, that is, a TTF and a DON moiety. Interestingly, the two 342+ and 372+ catenanes, containing four different donor and acceptor units, are fully desymmetrized. In all the catenanes, the electron donor ring surrounds the better electron acceptor BPY2+ unit, and in the case of catenanes 342+ and 372+, the better electron donor TTF unit is located inside the electron acceptor ring. Such conformations can be switched, altering the redox state of the donors and acceptors incorporated in the structure, as evidenced by the rich and complex electrochemical patterns exhibited by these catenanes. On the reduction side, the cyclic voltammogram of 322+ reveals four reversible monoelectronic processes. Spectroelectrochemical experiments and comparison with suitable model compounds [84,89] indicate that (1) the first process is consistent with the reduction of a BPY2+ unit engaged in CT interactions inside the cavity of the donor ring; (2) the second reduction concerns the monoreduced BPY+ unit, likely still inside the cavity of the donor ring; (3) the third process corresponds to the reduction of the inside PMI unit (the translocation probably occurs upon reduction of the BPY+ unit to its neutral form); and finally, (4) the fourth reduction involves PMI−, on the position of which relatively to the donor ring there is little information. In catenanes 332+ and 342+, because of the stronger electron-donating power of the crown ethers comprised in the structure, the BPY2+ unit becomes more difficult to be reduced than in catenane 322+ so that its second reduction overlaps with the first reduction of the PMI unit (Figure 12.24). The reduction pattern of 332+ and 342+ (Figure 12.24a) consists therefore of three reversible waves: (1) the first process concerns the monoelectronic reduction of the BPY2+ unit located inside the donor ring as confirmed by spectroelectrochemical experiments (for 342+, Figure 12.24b); (2) the second process, which involves the exchange of two electrons, can be unambiguously assigned to the overlapping reduction of the inside BPY+ and the alongside PMI [90]; and (3) the third process, which is monoelectronic, corresponds to the PMI−/2− reduction. The conformation changes coupled with the four-electron reduction processes for the 322+–342+ PMI-containing catenanes are summarized in Figure 12.25.

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Organic Electrochemistry

O

O

O

O

O

+N

C

B

A

322+–372+

+N

O

O

A

O

O

322+

N–

–N PMI

O

DOB

DON

O

O

332+

N–

–N O

O

O

O S

–N

N–

O

O

O

O

–N

–N

352+

O

362+

N– O

O

O

O

342+

O NDI

O –N

S

S S TTF

N–

O O

[2]Catenane

C

B

O

O

O

O

N–

S

S

S

S

372+

O

FIgURE 12.23 General structure and component units of the desymmetrized catenanes 322+–372+.

In the case of the three 35 2+ –372+ NDI-based catenanes, the reduction pattern comprises four reversible monoelectronic processes regardless of the nature of the donor rings. This behavior can be explained considering that the NDI unit is easier to be reduced than PMI. As a consequence, the processes observed can be assigned to BPY2+/+, NDI0/−, BPY+/0, and NDI−/2−, respectively, in the order of increasing reduction potential. The proposed conformational changes coupled with these four one-electron reduction processes are as follows (Figure 12.26): (1) the electron donor ring remains around the BPY+ radical cation after the first reduction; (2) the second reduction is indicated by spectroelectrochemical experiment to be the NDI0/− process [90], and the potential value, more consistent with the reduction of the NDI moiety in the alongside position, suggests that the donor ring still encircles the monoreduced BPY+ unit; (3) the remarkably negative potential value found for the third process, assigned to the BPY+/0 reduction, is in agreement with the assumption that no translocation of the donor component has occurred; and (4) the fourth process is attributed to the reduction of

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Electrochemically Driven Supramolecular Devices 6

5 1, 4

2

3

5 μA

–2.0

–1.5

–1.0

(a)

–0.5

+0.5

+1.0

E (V vs SCE)

0.3 A 0.2

1

0 V vs Ag

2

–0.60 V

3

–0.90 V

4

0 V vs Ag

5

+0.60 V

6

+1.00 V

A 0.1

0.1

0

0 400 (b)

600

800 λ (nm)

400

1000 (c)

600

800 λ (nm)

1000

FIgURE 12.24 Voltammetric and spectroelectrochemical response of catenane 342+ in CH3CN at room temperature. (a) Cyclic voltammetric curve (conditions: 0.49  mM, tetraethylammonium hexafluorophosphate 73 mM as supporting electrolyte, 200 mV/s, glassy carbon working electrode). (b) Absorption spectra observed before (full line) and after exhaustive reduction at −0.60 V (dashed line) and −0.90 V (dotted line) versus an Ag quasireference electrode. (c) Absorption spectra observed before (full line) and after exhaustive oxidation at +0.60 V (dashed line) and +1.00 V (dotted line) versus an Ag quasireference electrode. The numbered arrows in (a) mark the potential values at which the corresponding curves in (b) and (c) were recorded in the spectroelectrochemical experiments. (From Ceroni, P., Credi, A., Venturi, M., and Schalley, C.A. eds.: Analytical Methods in Supramolecular Chemistry. pp. 371–457. 2012. Copyright Wiley-VCH Verlag GmbH & Co. KGaA. Reproduced with permission.)

the NDI− unit, on the position of which relatively to the donor ring, there is little information, as noted previously for the PMI-based catenanes. On the oxidation side, only the behavior of the fully desymmetrized catenanes 342+ and 372+ is discussed because they are particularly interesting from the viewpoint of molecular machines. Their electrochemical patterns are very similar and consist of three oxidative processes (for 342+, Figure 12.24a): the first two (Figure 12.24c) are assigned to the two consecutive monoelectronic TTF oxidations [61], while the third one is ascribed to the oxidation of the DON unit.

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BPY reduction

+

+1e–

+

–1e–

+

+1e– –1e– –1e–

PMI reduction

–2e–

+1e–

+2e– –

=

DOB–DOB

322+

DON–DON

332+

TTF–DON

342+

–1e–

+1e–

– –

FIgURE 12.25 (See color insert.) The conformational changes associated with the four reduction processes for catenanes 322+–342+. Horizontal and vertical processes represent the BPY-centered and PMI-centered reductions, respectively. In the upper part of the scheme, gray arrows refer to the behavior of 322+, while black arrows describe the behavior of 332+ and 342+. The dotted ellipses indicate that the interactions are turned off, and there is little information about the donor ring location.

The first and second TTF oxidations exhibit the same features observed for the previously studied catenane 314+ [58b,86] and can be interpreted as follows: after the TTF0/+ oxidation, the electron donor ring circumrotates with respect to the electron-accepting ring, delivering the DON unit into its cavity. This electrochemical study evidences that in the desymmetrized 322+–372+ catenanes, the inside/ alongside topological preference, dominated by the intercomponent CT interactions, can be modulated reversibly upon reduction of the electron acceptors or oxidation of the electron donors in the cases of fully desymmetrized 342+ and 372+ catenanes. Such a feature makes these catenanes appealing structures for the construction of molecular machines and, in a perspective, rotary motors. It  should also be noted that upon electrochemical stimulation in a relatively narrow and easily accessible potential window, these interlocked molecules can be reversibly switched among several (six and seven for 342+ and 372+, respectively) states, all characterized by distinct electronic and optical properties. Such a possibility could open interesting routes for the development of molecular electronic devices that go beyond binary logic. Recently, a redox-switchable catenane composed of two interlocking 114+ macrocycles has been reported, and the electrochemical switching between six experimentally accessible redox states from within the total of nine states has been demonstrated [63]. A catenane composed of two identical benzylic amide macrocycles was also investigated by cyclic voltammetry [91]. Computer simulation of the voltammetric data, together with quantum chemical calculations, suggests that reduction of the macrocycles is followed by their irreversible

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Electrochemically Driven Supramolecular Devices

BPY reduction

+

+1e–

+

–1e–

+

NDI reduction

–1e–

+1e– –

+

–

+1e– –1e–

–1e–

=

+1e–

2+

DOB–DOB

35

DON–DON

362+

TTF–DON

372+

– –

FIgURE 12.26 (See color insert.) The conformational changes associated with the four reduction processes for catenanes 352+–372+. Horizontal and vertical processes represent the BPY-centered and NDI-centered reductions, respectively. The dotted ellipses indicate that the interactions are turned off, and there is little information about the donor ring location.

chemical soldering, owing to the formation of a C–C bond between two reduced carbonyl groups. Hence, the electrochemical stimulus can be used to prevent the mutual rotation of the two rings, although the irreversibility of the reaction limits further developments. Other examples of electrochemically driven switching processes concern metal-based catenanes [92–97], including heterodinuclear bis-macrocyclic transition-metal complexes [98]. A representative case is catenane [38•Cu]+ (Figure 12.27) that incorporates two identical macrocyclic components comprising terpyridine and phenanthroline ligands. The Cu(I) ion is coordinated tetrahedrally by the two phenanthroline ligands, whereas the two terpyridine ligands are located well away from each other [97]. The cyclic voltammogram of [38•Cu]+ contains a reversible wave (+0.63 V relative to the SCE) associated with the tetracoordinated Cu(II)/Cu(I) redox couple. The visible absorption spectrum of the catenane contains a metal-to-ligand charge-transfer band at 439 nm for the tetracoordinated Cu(I) chromophore. On electrochemical oxidation of [38•Cu]+ or on treatment with NOBF4, the tetracoordinated Cu(I) center is converted into a tetracoordinated Cu(II) ion that has an absorption band at 670 nm. The intensity of this band decreases with time, however. Indeed, in response to the preference of the Cu(II) ion for a coordination number higher than four, one of the two macrocycles circumrotates through the cavity of the other, affording a pentacoordinated Cu(II) ion. Subsequently, the other macrocycle undergoes a similar circumrotational process, yielding a hexacoordinated Cu(II) ion, which gives, instead, a weak absorption band at 687 nm. Electrolysis (−1.0 V) of the acetonitrile solution of the catenane reduces the hexacoordinated Cu(II) center back to a hexacoordinated Cu(I) ion. In response to the preference of Cu(I) for a tetracoordinated geometry, the two macrocycles circumrotate through the cavity of each other in turn, affording the original conformation quantitatively.

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Organic Electrochemistry O N

+

O N

N

N CuI

N N

N

O

–e–

CuII

N N

N

O

[38 Cu]+

CuI

CuI

CuII

+e–

CuII

[38 Cu]2+

FIgURE 12.27 Circumrotation of the macrocyclic components of catenane [38•Cu]+ controlled reversibly by oxidizing–reducing the metal center.

D.

MOLECULAR MACHINES WORKING ON SURFACES

As already noticed, the design, synthesis, and operation of multicomponent molecular systems capable of performing specific, directional mechanical movements under the action of a defined energy input—namely, molecular machines—constitute a fascinating challenge in the field of nanoscience. Most of these studies have been performed in solution where a huge number of molecules behave independently from one another because they cannot be addressed individually and hence controlled. Such an incoherent behavior is the major impediment to designing and realizing systems capable of performing useful functions. It seems, therefore, reasonable that for several real applications, the molecular machines have to be interfaced with the macroscopic world by ordering them in some way so that they can work coherently and can be addressed in space. Viable possibilities include deposition on surfaces, incorporation into polymers, organization at interfaces, or immobilization into membranes or porous materials [99–107]. 1. Molecular Machines Immobilized on Electrodes The formation of SAMs on metallic gold is an effective yet simple method to attach molecules to surfaces [24]. Rotaxanes with suitable functionalization have been extensively employed to prepare

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467

Electrochemically Driven Supramolecular Devices Fast e–

+N

Au

N

+

N

S

Au

N

N

S

N +N

+2e–

N+

Fe O NH O N N H N

H N

O N H

O Trans-40

394+

–2e–

Vis

UV

e– Slow Fe O

N Au S

NH

N+ H O N

N N

N

O

N +N

O N H

Au S

N

N H

N

N cis-40

(a)

(b)

FIgURE 12.28 (a) The electrochemically driven ring shuttling in rotaxane 394+ incorporated into an Au-SAM. (b) The surface-bound photoswitchable rotaxane 40 capable of transducing an optical signal into an electronic signal by means of the photocontrolled ring shuttling.

Au-SAMs [108,109]. The first studies on mechanical shuttling in a SAM, however, were carried out on systems in which the molecular components became mechanically interlocked as a consequence of surface immobilization [110]. In this setting, the metal surface plays the dual role of a stopper and an interface (electrode). A monolayer of the rotaxane 394+ (Figure 12.28a), consisting of the electron-accepting cyclophane 94+ threaded on a molecular axle, which includes an electron-donating di-iminobenzene unit and is stoppered by an adamantane moiety, was assembled on a gold electrode [110b]. The tetracationic ring, which is originally located on the di-iminobenzene unit by virtue of electron donor–acceptor interactions, is displaced toward the electrode upon one-electron reduction of its two bipyridinium units at −0.53 V versus the SCE, owing to disruption of the donor–acceptor interactions and electrostatic attraction to the electrode (Figure 12.28a). Reoxidation of the reduced cyclophane at −0.33 V versus SCE causes ring shuttling to the original di-iminobenzene site. The position of the tetracationic and dicationic (reduced) cyclophane rings and the shuttling rate constants (80 s−1 and 320 s−1 at 298 K for reduction- and reoxidation-induced processes, respectively) were determined by chronoamperometry and impedance measurements. Investigation of the temperature dependence of the shuttling rates showed [110c] that the reduction-induced shuttling is an energetically downhill process with no measurable activation barrier, whereas reoxidation-induced shuttling requires thermal activation. The lack of an energy barrier in the former case is in agreement with the fact that the shuttling is mainly driven by coulombic attraction of the still positively charged cyclophane toward the negatively polarized electrode. Therefore, ring shuttling does not need assistance by thermal energy, in contrast to the typical operation of molecular motors. The surface-bound rotaxane trans-40 (Figure 12.28b) consists of a ferrocene-functionalized β-CD macrocycle threaded on a molecule containing a photoisomerizable azobenzene unity and

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a long alkyl chain [110a]. A monolayer of trans-40 was self-assembled on a gold electrode. The azobenzene unit in the trans configuration is complexed by β-CD; photoisomerization to the cis form renders complexation sterically impossible, so that the β-CD ring is displaced to the alkyl component. Back-photoisomerization restores the original trans configuration. The position of the β-CD-tethered ferrocene unit was determined by chronoamperometry. A fast current decay (k = 65 s−1) was observed for the trans isomer, implying that the ring component is close to the electrode surface (Figure 12.28b). Photoisomerization of the monolayer to the cis state resulted in a chronoamperometric transient characterized by a substantially lower electron transfer rate (k = 15 s−1). This result indicates that in cis-40, the β-CD ring is more distant from the electrode surface. Owing to the reversibility of azobenzene photoisomerization, a cyclic pattern for the rate constant of the heterogeneous electron transfer process was observed. In this optoelectronic system, optical information is transduced by a mechanical shuttling to an electronic signal [110,111]. Another clever approach involves the immobilization of bistable molecular shuttles containing a pyridine residue in the macrocyclic component onto an Au-SAM of 11-mercaptoundecanoic acid [112]. The grafting is achieved because of the formation of hydrogen bonds between the pyridine unit of the rotaxane molecules and the carboxylic groups at the top of the SAM. Electrochemically driven shuttling in one of these systems was shown to occur on the millisecond timescale by chronoamperometry [112b]. Surface-mounted molecular rotary motors are also extremely interesting, both from a basic viewpoint [113] and because they could find applications in a variety of molecular-size devices and machines, for example, in the fields of nanoelectronic, nanophotonics, and nanofluidics [114]. A family of molecular rotors (e.g., compound 41 in Figure 12.29a) has been designed to perform rotation under electrochemical stimulation [115–117]. The molecules have a piano-stool structure with a stator meant to be grafted on an oxide surface and a rotor bearing redox-active groups, so that addressing the molecule with nanoelectrodes would trigger rotation (Figure 12.29b). To avoid intramolecular electron transfer between two electroactive units, which would compete with rotation, insulating spacers based on platinum acetylide units were inserted into the structure, but several difficulties remain to be overcome. 2. Solid-State Electronic Circuits The redox-switching behavior observed for solid-supported thin films of bistable catenanes [118] and rotaxanes [108ab,109a] encouraged attempts to incorporate such molecules in electrically addressable solid-state devices [119]. A Langmuir monolayer of the TTF–DNP rotaxane 424+ (Figure 12.30a) was transferred onto a photolithographically patterned polycrystalline silicon electrode [120]. The patterning was such that the film was deposited along several parallel lines of poly-Si on the electrode. A second set of orthogonally oriented wires was then deposited on top of the first set such that a crossbar architecture is obtained. This second set of electrodes consisted of a 5 nm thick layer of Ti, followed by a 100 nm thick layer of Al. By this approach, an array of junctions, each one addressable individually, was constructed (Figure 12.30b). In the first setup, the wire electrodes were a few micrometers wide, but the scalability of the fabrication method allowed the construction of wires less than 100 nm in width, yielding junctions with areas of 0.005–0.01 μm2 and containing about 5000 rotaxane molecules [119]. The mechanism for conduction is by electron tunneling through the single-molecule thick layer between the junction electrodes. Thus, any change in the electronic characteristics of the interelectrode medium is expected to affect the tunneling efficiency and change the resistance of the junction. It should be noticed that such devices are conductors, not capacitors. Experiments were carried out by applying a series of voltage pulses (between +2.0 and −2.0 V) and reading, after each pulse, the current through the device at a small voltage (between +0.2 and −0.2 V) that does not

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Electrochemically Driven Supramolecular Devices

Fe

Et3P–Pt–PEt3 Fe

Fe Et3P Pt PEt3

PEt3 Pt Et3P

Et3P

PEt3 Pt

Fe

Pt PEt3

RuII

N N

N

Et3P

Fe

N N

N B O

O

O

H

O O

(a)

O

41

e– +

–

+

1

5

5 4

–

+ 1 2

2 3

4

3

e– +

+ 5

–

+

e– 5

1

+ 1

–

4

4 3

2

3

2

(b)

FIgURE 12.29 (a) Molecular rotor 41 designed to perform rotation under electrochemical stimulation in a nanoelectrode junction. (b) Schematic representation of the proposed operation mechanism. A potential difference applied across the nanojunction results in oxidation of the ferrocene group nearest the anode, which is then electrostatically repelled away toward the cathode. The oxidized unit close to the cathode is reduced, while a new ferrocene unit is oxidized at the anode. Overall, the passage of each electron across the nanojunction results in the clockwise rotation of the upper part of the molecule by one-fifth of a turn. (From Balzani, V., Credi, A., and Venturi, M.: Molecular Devices and Machines, 2nd edn. p. 492. 2008. Copyright Wiley-VCH Verlag GmbH & Co. KGaA. Reproduced with permission.)

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Organic Electrochemistry 424+ +

O

N+

N

O

S

O

S

O

S S TTF +N

OCH3 O

O O

O

O

O

O

OCH3

O

O DNP

O

O

OCH3

N+

(a)

Ti-Al

+

Poly-Si (b)

FIgURE 12.30 (See color insert.) Structure formula of bistable rotaxane 424+ (a) and schematic representation of solid-state junctions consisting of a monolayer of 424+ sandwiched between poly-Si and Ti-Al crossbar electrodes (b). (From Silvi, S., Venturi, M., Credi, A., J. Mater. Chem., 19, 2279–2294, 2009. Reproduced by permission of The Royal Society of Chemistry.)

affect switching. The current (read)–voltage (write) curve displays a highly hysteretic profile, making the rotaxane junction device interesting for potential use in RAM storage. The current–voltage curve was interpreted on the basis of the mechanism illustrated in Figure 12.31, which is derived from the behavior of the same rotaxane in solution [87]. Coconformation A is the switch open state and coconformation D the switch closed state of the device. When the TTF unit of 424+ is oxidized (+2 V, state B), a coulombic repulsion inside the tetracationic cyclophane component is generated, which causes the displacement of the latter and formation of state C in which the ring encircles the DNP unit (note that, in solution at +2 V vs SCE, TTF undergoes two-electron oxidation and DNP is also oxidized [121]). When the voltage is reduced to near-zero bias, a metastable state D is obtained, which, however, does not return to state A. The initial coconformation can in fact be restored only via states E and F in which the bipyridinium units of the cyclophane component are reduced (in solution, at the potential value used, −2 V, each bipyridinium unit undergoes twoelectron reduction [121]). Most likely, the reduction of the bipyridinium units weakens the chargetransfer interaction with the DNP unit, thereby decreasing the barrier that hinders the replacement of the cyclophane on the TTF site. An analogous mechanism was used to interpret the behavior of solid-state devices containing other TTF-based bistable interlocked molecules [119–122]. The metastable state corresponding to coconformation D (Figure 12.31) was in fact observed for a number of different bistable rotaxanes and catenanes in a variety of environments (solution, SAM, and solid-state polymer matrix) [123]. More recently, by the use of the same paradigms and the same bistable rotaxane 424+ described earlier, a molecular electronic memory with an amazingly high density of 1011 bits cm−2 was constructed by sandwiching a monolayer of the rotaxane between arrays of nanoelectrodes in a crossbar arrangement [124]. The realization of this device relies on a novel method for producing ultradense, highly aligned arrays and crossbars of metal or semiconductor nanowires with high aspect ratios [125]. It was estimated that each junction acting as a memory element consists of approximately 100 rotaxane molecules. For practical reasons, only 128 (16 × 8 contacts) of the 160,000 memory cells

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471

Electrochemically Driven Supramolecular Devices Ti-Al

–2e

–2e

Poly-Si

F

A

+2e

E –2

B

+2e

D 0 E (V)

C +2

FIgURE 12.31 (See color insert.) Operation mechanism of solid-state molecular electronic switching devices based on bistable molecular shuttles like 424+. (From Silvi, S., Venturi, M., Credi, A., J. Mater. Chem., 19, 2279–2294, 2009. Reproduced by permission of The Royal Society of Chemistry.)

(400 × 400 nanowires) contained in the circuit were tested [124]. The measurements showed that 25% of the tested cells displayed good and reproducible switching, whereas 35% failed because of bad contacts or shorts, and the remaining 40% showed poor switching. This work is a compelling demonstration that the combination of top-down and bottom-up nanofabrication methods can lead to outstanding technological achievements. However, several aspects—such as stability, reliability, and ease of fabrication—need to be optimized before these systems can find real industrial applications [126]. 3. Electrochemically Induced Shuttling in Single Rotaxane Molecules Experiments on individual motor protein molecules have been crucial to understand their working mechanism [127]. The control and observation of motion of single artificial molecular machines is indeed a very challenging and highly stimulating task. In fact, while detailed kinetic and thermodynamic information on the machine and its working mechanism can be determined from ensemble experiments, the functionality as a nanoscale device can only be investigated by operating the system at the single molecule level. Although fluorescence microscopy can be useful in some cases [128], this kind of studies requires the use of scanning probe microscopy (SPM) techniques [129,130]. Several investigations have shown that movements of single molecules deposited on surfaces can be induced and observed by using an SPM tip. Among targeted systems are molecular rotors [113,131], polyrotaxanes [132], catenanes [133], wheelbarrows [134], rack-and-pinion devices [135], gears [136], and nanovehicles [137]. In a remarkable investigation, bistable rotaxane molecules similar to those described in Section IV.D.2, behaving as redox-controlled molecular shuttles in solution, were anchored laterally on gold surfaces and investigated using an electrochemical scanning tunneling microscope (STM) [138]. The rotaxane molecules (43 4+ in Figure 12.32) were bound at each end to a Au{111} surface by means of the disulfide groups attached to the stoppers on the dumbbell termini.

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Organic Electrochemistry 434+

+

+

N

N O O

O

O

O

O

S

S

S

S

O

O

O

O

O

O

O

O

O

O S S

+N

S

N+

S

–e– + +

+

+e–

+ +

FIgURE 12.32 (See color insert.) Structure of the bistable rotaxane molecule 434+ adsorbed on Au{111} and schematic representation of its redox-induced shuttling motion investigated using electrochemical scanning tunneling microscopy.

The molecules were assembled with a surface coverage of 5–6 molecules/1000 Å2 in orientations conducive to direct STM measurements of their station changes. The molecules were imaged at two different electrode potentials, +0.12 and +0.53 V versus Ag/AgCl, for which the TTF station is in the reduced (neutral) state and in its oxidized (cationic) state, respectively. A displacement of the tetracationic ring was observed and correlated with the redox states of the TTF station. The displacement exhibits partial reversibility upon stepping the potential back to +0.12 V. The trajectories of the rings determined from imaging of many such molecules suggest that the motion of the ring relative to the dumbbell in a surface-adsorbed molecule is affected by its local environment and the flexibility of the molecule [138]. It is envisaged that rotaxanes with rigid dumbbell components should enable a better visualization and exhibit consistent and reversible motion. 4. Electrically Driven Directional Motion of a Single Molecular Machine Single molecules that can move in a controlled manner on a nonmodified surface in response to external stimuli have to use chemical, electrical, or light energy to modulate their interaction with the surface in a way that generates motion. This extremely challenging objective was recently achieved with a nanocar molecule 44 (Figure 12.33a) that comprises four rotary motor units as the wheels [139]. Such units undergo unidirectional rotation of the fluorene moieties around the C=C double bond as a result of sequential configurational and conformational switching processes induced by electronic (e.g., light) and vibrational (e.g., thermal) excitation, respectively [140]. The direction of rotation of each motor unit is dictated by its chirality. As it is known that conformational and configurational switching of molecules on surfaces can be obtained by STM-induced vibrational and electronic excitations, it was envisaged that an STM tip could bring about the rotary motion of the motor units of 44 and, at the same time, afford single molecule imaging. After deposition onto a Cu{111} surface by sublimation, individual 44 molecules were imaged with an STM at 7 K. Imaging conditions are sufficiently mild that no changes are induced upon continuous scanning, whereas electronic and vibrational excitations of the molecule are induced upon application of appropriate voltage pulses to the STM tip. A translational motion of the molecule across the surface should be observed if the four motors all rotate in the same direction, a requirement that is only met by the meso-(R,S-R,S) isomer sketched in Figure 12.33a. This is indeed what was observed in the experiments [139]. Additional support for the electrically driven directional motion comes from investigations aimed at probing the effect of the chirality and geometry of the molecule on its motion. For example, two

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Electrochemically Driven Supramolecular Devices

R

R N

N

S

S meso-(R,S-R,S)-Y

(a)

meso-(R,S-R,S) ''correct landing''

meso-(R,S-R,S) ''wrong landing''

(R,R-R,R) enantiomer

(b)

FIgURE 12.33 (See color insert.) (a) Structure and cartoon representation of the meso-(R,S-R,S) isomer of the four-wheeled molecule 44. (b) Schematics of the directionality of the motion induced by the concerted rotation of the motor units. The arrows indicate the direction in which the rotary action of the individual motor units propels the molecule on the surface. Two distinct landing geometries of the meso-isomer lead either to directional movement (left) or to no movement at all (center), whereas (R,R-R,R) (right) or (S,S-S,S) enantiomers move randomly.

geometries must be considered for the meso-isomer because free rotation around the bis-alkyne C–C single bond of the frame (gray arrow in Figure 12.33a; pink in the color insert) is locked upon adsorption (Figure 12.33b). When the meso-isomer is adsorbed on the surface in the proper orientation (correct landing), conrotatory motion of the four motor units moves the molecule along. Conversely, in the wrongly landed meso-isomer, the effects arising from the movement of the motor units cancel out, thereby precluding translation. The individual (R,R-R,R) or (S,S-S,S) enantiomers of 44 transferred on the surface from the racemic mixture of the compound were found to spin and randomly move across the surface [139]. This behavior can be explained considering that the motor units on opposite sides of the molecule rotate in a disrotatory fashion, ideally causing the molecules to spin; in a nonideal case, the molecules exhibit random translational motion in addition to the spinning motion. These pioneering experiments demonstrate that a single molecule with intrinsic motor functions is capable of converting an external energy input into directional motion along a surface and that the type of movement is a direct consequence of the molecular design.

V. CONCLUDINg REMARKS In this chapter, we showed that redox active complex systems, as the supramolecular devices here described, can be fully characterized by using the various kinds of electrochemical techniques. They provide indeed a fingerprint of the analyzed systems giving fundamental information on (1) the spatial organization of the redox sites within the molecular and supramolecular structure, (2) the entity of the interactions between such sites, and (3) the kinetic and thermodynamic stabilities of the reduced/oxidized and charge-separated species. Electrochemistry is, therefore, a powerful tool to read the state of the system.

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A further aim of this chapter is to evidence that in suitable designed systems, electrochemistry can play a more important role. By causing the occurrence of endoergonic electron transfer processes, electrochemistry can, indeed, provide the energy needed to modify the noncovalent interactions that stabilize a certain structure promoting mechanical movements. In such cases, electrochemistry plays the dual role of writing and reading a system: by means of electrons and/ or holes, it supplies the energy to make these systems work as molecular machines, and by means of the various electrochemical techniques, it is used for controlling and monitoring the operation performed by the system. In this regard, selected examples of molecular machines have been described. They show that, although investigations in solution are fundamental to understand the operation mechanisms of such systems, an important step for any kind of applications in technology requires that they are interfaced with the macroscopic world and ordered in some way so that they can behave coherently and can be addressed in space. Another important challenge is the control and observation of the motion of single artificial machines. The recent achievements obtained in these directions, some of which have been reviewed here, seem to indicate that useful materials and devices based on artificial molecular machines will see the light in a not too distant future. The systems described here evidence that electrochemists have learned how to deal with increasingly complex molecular and supramolecular structures. However, it must be noticed that electrochemistry is only a part of the game. As the complexity of the systems studied increases, the contribution from many disciplines in a joint and collaborative effort is needed. The goal of transforming molecular devices and machines into practically useful products requires, indeed, that people belonging to different fields, like chemistry, solid-state physics, biology, computer science, mathematics, materials sciences, etc., work together and learn a common language.

ACKNOWLEDgMENTS Financial support from Ministero dell’Università e della Ricerca (PRIN 2010CX2TLM "InfoChem" and FIRB 2010RBAP11C58Y "Nanosolar") and Ministero degli Affari Esteri e Cooperazione Internazionale (Progetto Italia-USA) is gratefully acknowledged.

REFERENCES 1. Lehn, J.-M. Supramolecular Chemistry: Concepts and Perspectives; Wiley-VCH: Weinheim, Germany, 1995. 2. Balzani, V.; Scandola, F. Supramolecular Photochemistry; Horwood: Chichester, U.K., 1991. 3. Marcaccio, M.; Paolucci, F.; Roffia, S. In Trends in Molecular Electrochemistry; Pombeiro, A. J. L.; Amatore, C. (eds); Dekker: New York, 2004, p. 223. 4. Joachim, C.; Launay, J. P. Nouv. J. Chem., 1984, 8, 723. 5. Balzani, V.; Moggi, L.; Scandola, F. In Supramolecular Photochemistry; Balzani, V. (ed.); Reidel: Dordrecth, the Netherlands, 1987, p. 1. 6. Lehn, J.-M. Angew. Chem. Int. Ed. Engl. 1990, 29, 1304. 7. Balzani, V.; Credi, A.; Venturi, M. Chem. Eur. J. 2002, 8, 5524. 8. Balzani, V.; Credi, A.; Venturi, M. Molecular Devices and Machines—Concepts and Perspectives for the Nanoworld, 2nd edn.; Wiley-VCH: Weinheim, Germany, 2008. 9. Bard, A. J.; Faulkner, L. R. Electrochemical Methods. Fundamentals and Applications, 2nd edn.; Wiley: Hoboken, NJ, 2001. 10. Kaifer, A. E.; Gómez-Kaifer, M. Supramolecular Electrochemistry; Wiley-VCH: Weinheim, Germany, 1999. 11. Balzani, V. (ed.). Electron Transfer in Chemistry; Wiley-VCH: Weinheim, Germany, 2001. 12. Hodes, G. (ed.). Electrochemistry of Nanomaterials; Wiley-VCH: Weinheim, Germany, 2001. 13. Willner, I.; Katz, E. (eds.). Bioelectronics; Wiley-VCH: Weinheim, Germany, 2005. 14. Ceroni, P.; Credi, A.; Venturi, M. (eds.). Electrochemistry of Functional Supramolecular Systems; Wiley: Hoboken, NJ, 2010.

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Electrochemically Driven Supramolecular Devices 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43. 44. 45. 46. 47.

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475

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110. (a) Willner, I.; Pardo-Yssar, V.; Katz, E.; Ranjit, K. T. J. Electroanal. Chem. 2001, 497, 172. (b) Katz, E.; Lioubashevsky, O.; Willner, I. J. Am. Chem. Soc. 2004, 126, 15520. (c) Katz, E.; Baron, R.; Willner, I.; Richke, N.; Levine, R. D. ChemPhysChem 2005, 6, 2179. 111. Shipway, A. N.; Katz, E.; Willner, I. Struct. Bond. 2001, 99, 237. 112. (a) Cecchet, F.; Rudolf, P.; Rapino, S.; Margotti, M.; Paolucci, F.; Baggerman, J.; Brouwer, A. M.; Kay, E. R.; Wong, J. K. Y.; Leigh, D. A. J. Phys. Chem. B 2004, 108, 15192. (b) Fioravanti, G.; Haraszkiewicz, N.; Kay, E. R.; Mendoza, S. M.; Bruno, C.; Marcaccio, M.; Wiering, P. G.; Paolucci, F.; Rudolf, P.; Brouwer, A. M.; Leigh, D. A. J. Am. Chem. Soc. 2008, 130, 2593. 113. Kottas, G. S.; Clarke, L. I.; Horinek, D.; Michl, J. Chem. Rev. 2005, 105, 1281. 114. Wang, B.; Král, P. Phys. Rev. Lett. 2007, 98, 266102. 115. Rapenne, G. Org. Biomol. Chem. 2005, 3, 1165. 116. Carella, A.; Coudret, C.; Guirado, G.; Rapenne, G.; Vives, G.; Launay, J.-P. Dalton Trans. 2007, 177. 117. Vives, G.; Gonzales, A.; Jaud, J.; Launay, J.-P.; Rapenne, G. Chem. Eur. J. 2007, 13, 5622. 118. Asakawa, M.; Higuchi, M.; Mattersteig, G.; Nakamura, T.; Pease, A. R.; Raymo, F. M.; Shimidzu, T.; Stoddart, J. F. Adv. Mater. 2000, 12, 1099, and references therein. 119. (a) Collier, C. P.; Mattersteig, G.; Wong, E. W.; Luo, Y.; Beverly, K.; Sampaio, J.; Raymo, F. M.; Stoddart, J. F.; Heath, J. R. Science 2000, 289, 1172. (b) Collier, C. P.; Jeppesen, J. O.; Luo, Y.; Perkins, J.; Wong, E. W.; Heath, J. R.; Stoddart, J. F. J. Am. Chem. Soc. 2001, 123, 12632. 120. (a) Luo, Y.; Collier, C. P.; Jeppesen, J. O.; Nielsen, K. A.; Delonno, E.; Ho, G.; Perkins, J.; Tseng, H.-R.; Yamamoto, T.; Stoddart, J. F.; Heath, J. R. ChemPhysChem 2002, 3, 519. 121. Tseng, H.-R.; Vignon, S. A.; Celestre, P. C.; Perkins, J.; Jeppesen, J. O.; Di Fabio, A.; Ballardini, R.; Gandolfi, M. T.; Venturi, M.; Balzani, V.; Stoddart, J. F. Chem. Eur. J. 2004, 10, 155. 122. (a) Pease, A. R.; Jeppesen, J. O.; Stoddart, J. F.; Luo, Y.; Collier, C. P.; Heath, J. R. Acc. Chem. Res. 2001, 34, 433. (b) Stewart, D. R.; Ohlberg, D. A. A.; Beck, P. A.; Chen, Y.; Williams, R. S.; Jeppesen, J. O.; Nielsen, K. A.; Stoddart, J. F. Nano Lett. 2004, 4, 133. 123. (a) Flood, A. H.; Peters, A. J.; Vignon, S. A.; Steuerman, D. W.; Tseng, H.-R.; Kang, S.; Heath, J. R.; Stoddart, J. F. Chem. Eur. J. 2004, 10, 6558. (b) Choi, J. W.; Flood, A. H.; Steuerman, D. W.; Nygaard, S.; Braunschweig, A. B.; Moonen, N. N. P.; Laursen, B. W.; Luo, Y.; Delonno, E.; Peters, A. J.; Jeppesen, J. O.; Xu, K.; Stoddart, J. F.; Heath, J. R. Chem. Eur. J. 2006, 12, 261. 124. Green, J. E.; Choi, J. W.; Boukai, A.; Bunimovich, Y.; Johnston-Halperin, E.; Delonno, E.; Luo, Y.; Sheriff, B. A.; Xu, K.; Shin, Y. S.; Tseng, H.-R.; Stoddart, J. F.; Heath, J. R. Nature 2007, 445, 414. 125. Melosh, N. A.; Boukai, A.; Diana, F.; Gerardot, B.; Badolato, A.; Petroff, P. M.; Heath, J. R. Science 2003, 300, 112. 126. Ball, P. Nature 2007, 445, 362. 127. See, e.g., pioneering work: (a) Svoboda, K.; Schmidt, C. F.; Schnapp, B. J.; Block, S. M. Nature 1993, 365, 721. (b) Finer, J. T.; Simmons, R. M.; Spudich, J. A. Nature 1994, 368, 113. (c) Noji, H.; Yasuda, R.; Yoshida, M.; Kinosita Jr., K. J. Nature 1997, 386, 299. 128. Nishimura, D.; Takashima, Y.; Aoki, H.; Takahashi, T.; Yamaguchi, H.; Ito, S.; Harada, A. Angew. Chem. Int. Ed. 2008, 47, 6077. 129. Grill, L. J. Phys. Cond. Matter 2008, 20, 053001. 130. Lussis, P.; Svaldo-Lanero, T.; Bertocco, A.; Fustin, C.-A.; Leigh, D. A.; Duwez, A.-S. Nat. Nanotech. 2011, 6, 553. 131. (a) Gimzewski, J. K.; Joachim, C. Science 1999, 283, 1683. (b) Wintjes, N.; Bonifazi, D.; Cheng, F.; Kiebele, A.; Stoehr, M.; Jung, T.; Spillmann, H.; Diederich, F. Angew. Chem. Int. Ed. 2007, 46, 4167. (c) Michl, J.; Sykes, C. H. ACS Nano 2009, 3, 1042. 132. Shigekawa, H.; Miyake, K.; Sumaoka, J.; Harada, A.; Komiyama, M. J. Am. Chem. Soc. 2000, 122, 5411. 133. Payer, D.; Rauschenbach, S.; Malinowski, N.; Konuma, M.; Virojanadara, C.; Starke, U.; DietrichBuchecker, C.; Collin, J.-P.; Sauvage, J.-P.; Lin, N.; Kern, K. J. Am. Chem. Soc. 2007, 129, 15662. 134. Grill, L.; Rieder, K.-H.; Moresco, F.; Jimenez-Bueno, G.; Wang, C.; Rapenne, G.; Joachim, C. Surf. Sci. 2005, 584, L153. 135. Chiaravallotti, F.; Gross, L.; Rieder, K.-H.; Stojkovic, S. M.; Gourdon, A.; Joachim, C.; Moresco, F. Nat. Mater. 2007, 6, 30. 136. Manzano, C.; Soe, W.-H.; Wong, H. S.; Ample, F.; Gourdon, A.; Chandrasekhar, N.; Joachim, C. Nat. Mater. 2009, 8, 576. 137. (a) Shirai, Y.; Morin, J.-F.; Sasaki, T.; Guerrero, J. M.; Tour, J. M. Chem. Soc. Rev. 2006, 35, 1043. (b) Shirai, Y.; Osgood, A. J.; Zhao, Y. M.; Yao, Y. X.; Saudan, L.; Yang, H. B.; Chiu, Y. H.; Alemany, L. B.; Sasaki, T.; Morin, J. F.; Guerrero, J. M.; Kelly, K. F.; Tour, J. M. J. Am. Chem. Soc. 2006, 128, 4854.

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138. 139. 140. 141. 142.

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(c) Sasaki, T.; Tour, J. M. Org. Lett. 2008, 10, 897. (d) Sasaki, T.; Osgood, A. J.; Alemany, L. B.; Kelly, K. F.; Tour, J. M. Nano Lett. 2008, 10, 229. (e) Vives, G.; Kang, J.; Kelly, K. F.; Tour, J. M. Nano Lett. 2009, 11, 5602. (f) Vives, G.; Kang, J. H.; Kelly, K. F.; Tour, J. M. Org. Lett. 2009, 11, 5602. Ye, T.; Kumar, A. S.; Saha, S.; Takami, T.; Huang, T. J.; Stoddart, J. F.; Weiss, P. S. ACS Nano 2010, 4, 3697. Kudernac, T.; Ruangsupapichat, N.; Parschau, M.; Maciá, B.; Katsonis, N.; Harutyunyan, S. R.; Ernst, K.-H.; Feringa, B. L. Nature 2011, 479, 208. (a) Koumura, N.; Zijlstra, R. W. J.; van Delden, R. A.; Harada, N.; Feringa, B. L. Nature 1999, 401, 152. (b) Vicario, J.; Walko, M.; Meetsma, A.; Feringa, B. L. J. Am. Chem. Soc. 2006, 128, 5127. Ceroni, P.; Credi, A.; Venturi, M.; Schalley, C. A. eds., Analytical Methods in Supramolecular Chemistry, Wiley-VCH: Weinheim, Germany, 2012, pp. 371–457. Silvi, S.; Venturi, M.; Credi, A. J. Mater. Chem. 2009, 19, 2279–2294.

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13

Proton-Coupled Electron Transfers Cyrille Costentin, Marc Robert, and Jean-Michel Savéant

CONTENTS I. Introduction ............................................................................................................................ 481 II. Modeling Electrochemical-Concerted Electron and Proton Transfer Reactions ................... 483 A. Thermodynamics ............................................................................................................ 483 B. Modeling Concerted Proton–Electron Transfer Kinetics ............................................... 485 III. Competition between Stepwise (EPT and PET) and Concerted (CPET) Pathways ............... 488 IV. Intrinsic Characteristics of CPET Processes .......................................................................... 494 A. Similarity and Differences between ET, CPET, and CDET ........................................... 494 B. Hydrogen-Bonded Systems ............................................................................................. 496 C. Water (in Water) as Proton Acceptor ..............................................................................500 V. CPET for Bond Activation ...................................................................................................... 505 A. Breaking Bonds with Protons and Electrons .................................................................. 505 B. Activation of Molecules .................................................................................................. 507 VI. Concluding Remarks .............................................................................................................. 507 References ...................................................................................................................................... 508

I. INTRODUCTION The coupling between electron and proton transfers has a long experimental and theoretical history in chemistry and biochemistry. Proton-coupled electron transfer (PCET) reactions also play a critical role in a wide range of biological processes, including enzyme reactions, photosynthesis, and respiration. PCET is employed here as a general term for reactions in which both an electron and a proton are transferred, either in two distinct steps or in a single step [1]. We term the latter mechanism concerted proton and electron transfer (CPET) [2]. Other terms have been used in the literature to describe the same mechanism: electron transfer–proton transfer (ETPT) [3], or electron–proton transfer (EPT) [4], or multiple site–electron proton transfer (MS–EPT) [5]. Reactions in which the electron and proton transfers occur between the same donor and acceptor, that is, hydrogen-atom transfer, are not considered here. We are mainly interested in electrochemical PCET reactions in which electrons are flowing into or from an electrode, while protons are transferred between an acid and a base. Among the various ways of injecting or removing electrons, molecular electrochemistry, through nondestructive techniques such as cyclic voltammetry, has proved very useful in characterizing electron transfers (ETs) and deciphering mechanisms where chemical reactions, for example, proton transfer, are associated with ET [6]. This approach possesses several advantages. In the context of PCET, separation of the ET (the electrode) and proton transfer sites, required to distinguished CPET reactions from H-atom transfers, is readily achieved. Additionally, changing the electrode potential is an easy way of varying the driving force of the reaction and the current is an online measure of the reaction kinetics. It is, however, important to note that other injection or removal modes of the electron through thermal homogeneous reaction or photoinduced reaction can be used and that 481

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482

Organic Electrochemistry

QA 2H2O

O2

S4

P680

P680+

13.4

P680

P680

S0

S1 S2 P680 P + 680

S3

PheoD1

P680+ P680

11.1

10.6 8.2

(a)

ChlD1

(b)

10.4

10.2 13.8 Tyrz

Ala 344 CP43–Glu354 1.7 2.5

Asp 170 2.4

4 2.5

2.1 3 2.3

2.2

Asp 61

2.4

1

D1 H190 Mimic

Glu 189

H

1.8 2.2

Asp 342

Glu 333

5.1

OEC

2.6 4.9 Ca 2.5 2

PD2

TyrZ

GIn 165

CP43–Arg 357

PD1

ChlD2

O His 332

N

But

His 337

tBu

(c)

(d)

SCHEME 13.1 (See color insert.) Schematic view of photosystem II. (a) Kok cycle. (b) Structure of the reaction center of photosystem II showing the TyrZ –ChlD1(P680)–PheoD1–QA donor–chromophore–acceptor system, electron transfer from tyrosine (TyrZ) being coupled to proton transfer from histidine D1 H190. OEC, oxygen evolving complex. (c) One proposed schematic view of the OEC Mn4Ca Ala, alanine; Arg, arginine; Asp, aspartate; Glu, glutamate; His, histidine. The numbers are the distances in angstroms. In the labeling scheme, amino acids in black are in the first coordination sphere and those beyond in gray. (d) Aminophenol mimicking the TyrZ –histidine couple.

comparison of results obtained from various methods can be useful for a complete understanding of the PCET processes. Likewise, the electrochemical approach can be used to investigate PCET reactions taking place in a homogeneous context. Photosystem II (PSII) is a typical example of this approach. As shown in Scheme 13.1, oxidation of tyrosine in PSII is part of the charge transfer pathway between the chromophore and the oxygen-evolving complex (OEC). Oxidation of the tyrosine residue is coupled to proton transfer to a nearby histidine. To study the possible mechanistic pathways of such PCET, electrochemical investigation of a mimic (see Scheme 13.1) has proved useful as detailed later on in this chapter. As depicted in Scheme 13.2, PCET may follow either stepwise or concerted mechanisms. Thermodynamic and kinetic characterization of these pathways is the first purpose task of this chapter. In the discussion, it will appear that two conditions are required to favor concerted process vs. stepwise pathways: a favorable thermodynamical situation corresponding to high-energy intermediates in stepwise pathways and intrinsic favorable parameters corresponding mainly to the requirement of short distances between the proton donor and acceptor. A large part of this chapter is thus devoted to illustrating the dichotomy of the stepwise vs. concerted competition for PCET reaction controlled by thermodynamic parameters. Then, the intrinsic parameters of concerted pathways will be analyzed with particular emphasis on the role of hydrogen bond (H bond), H-bond relays, and ultimately peculiar behavior of water (in water) as proton acceptor. This chapter ends with a discussion on the implication of CPET in the activation of molecules. Although coupling of proton transfer to ET has a long history in organic electrochemistry involving

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483

Proton-Coupled Electron Transfers EPT –e– XR H + B

X OH + B –e– –

+e

CPET –e– R

+

X + HB

XO + HB+

+e+

PET

SCHEME 13.2

One-electron/one-proton PCET mechanistic scheme.

in particular carbon acid–base reactions [6], the examples discussed in this chapter are related to oxygen or nitrogen acid–base due to their involvement in biological systems [7] or small molecules activation in the context of contemporary energy challenges [8].

II.

MODELINg ELECTROCHEMICAL-CONCERTED ELECTRON AND PROTON TRANSFER REACTIONS

A.

THERMODYNAMICS

The global equation for a one-electron/one-proton PCET reaction (Scheme 13.2) reads X R H + B ⇌ X O + e − + BH + It is thermodynamically characterized by a standard potential, EX0 O + BH+ /XR H+B =

µ 0XO + µ 0BH+ − µ 0XR H − µ 0B F

(the μ0 are the standard chemical potentials of the subscript species) shown on the oblique straight line of the Pourbaix diagram of Figure 13.1, which relates the equilibrium potential when [XRH] = [XOH], or apparent standard potential, Eap0 , to the pH of the solution Eap0 = EX0 O ,H+ /XR H −

RT ln 10 pH F

(13.1)

In this representation, the nature of the acid–base couples involved does not matter insofar the reactions are all at equilibrium. The standard potential involved in Equation 13.1 may thus be equated to EX0 O ,H+ /XR H =

µ 0XO + µ H0 + − µ 0XR H , F

which does not refer to any particular acid–base couple and in which µ 0H+ is the standard chemical potential of the proton in the solvent under consideration whatever the structure of the solvated proton. Besides the oblique line, the Pourbaix diagram also shows two horizontal lines corresponding to the standard potential of the protonated and deprotonated redox couples: Eap0 = EX0 OH/XR H and Eap0 = EX0 O /XR , respectively.

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484

Organic Electrochemistry E0ap E0EPT

XOH XO

E0CEPT XRH E0PET

XR pK

XOH

pH

pK

XRH

pKHB (= 0 if HB = H2O)

FIgURE 13.1 Thermodynamics of PCET reactions (Scheme 13.2) and Pourbaix diagram.

Overall, the Pourbaix diagram (Figure 13.1) provides a map of the zones of thermodynamic stability of the various species involved. The standard potential corresponding to a specific acid– base couple may thus be derived from the Pourbaix diagram by formally equating the pH to the pK of this acid–base couple, as pictured in Figure 13.1. For, for example, water, which is the proton acceptor in a number of cases, the pertinent standard potential is the value of Eap0 at pH 0 and not a “standard potential” that would depend on the pH of the aqueous solution. In the earlier mentioned thermodynamic analysis, we have implicitly assumed that the activity coefficient of all intervening species is equal to unity. If not the case, the activity coefficient, γ, should be introduced in the standard chemical potential by replacing in the aforementioned equations, μ0 by μ0 + RT ln(γ). With reference to Scheme 13.2 and Figure 13.1, reactions going from XRH into XO and reverse may follow stepwise pathways (EPT for ET followed by proton transfer or PET for proton transfer followed by ET) and thus requires the intermediacy of XOH or XR intermediates. Alternatively or competitively, a concerted mechanism involving an acid–base couple HB/B as proton donor/acceptor couple, characterized by pKHB, may take place, thus skipping these intermediates. Distinction and competition between these mechanisms rest on kinetics, but the thermodynamic framework provided by the driving force characterizing each pathway is an essential requisite (the driving force of a reaction is here precisely defined as the opposite of the standard free energy of this reaction, −ΔG 0). The two stepwise reaction pathways are governed by two driving forces, one for ET and one for proton transfer, whereas there is a single driving force for the CPET pathway: EPT :

(

)

0 0 = RT ln 10 ( pK HB − pK XOH ) −∆GET = F E − EX0 OH/XR H , − ∆GPT

PET :

(

0 −∆GP0T = RT ln 10 ( pK HB − pK XR H ) , − ∆GET = F E − EX0 O /XR

)

CPET :

(

0 −∆GCPET = F E − EX0 O +HB/XR H + B

)

(E is the electrode potential, or for homogeneous ETs, the standard potential of the redox couple that provides or receives the electrons to or from the system). It is important to emphasize that, for a given proton donor/acceptor couple HB/B, the CPET driving force does not depend on pH [9–11]. Results on the oxidation of tyrosine and tryptophan in water have been interpreted as resulting from a CPET reaction with water as the proton endowed with a pH-dependent driving force, in the framework of a brute-force application of the Marcus theory for outer sphere ET (see Figure 13.4

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485

Proton-Coupled Electron Transfers XRH + B

SCHEME 13.3

XRH ... B

XO ...HB+

XO + HB+

Hydrogen bonding in CPET pathway.

in Reference 11a and Figure 13.5 in Reference 11b). This notion of pH-dependent driving force is nothing else than contradictory to the second law of thermodynamics. It should also be noted that CPET reactions may involve hydrogen bonding as sketched in Scheme 13.3. Examples will be given in Section IV.B. Kinetic analysis of the stepwise pathways (EPT and PET) follows the classical treatments of EC and CE reaction schemes, in, for example, cyclic voltammetry where a precise description of the various possible cases (kinetic control by ET and/or chemical reaction) is available [6]. In contrast, CPET processes require modeling of the kinetics, leading to the formulation of a rate law as discussed in Section II.B.

B.

MODELING CONCERTED PROTON–ELECTRON TRANSFER KINETICS

The following equation is a general expression of the rate law that relates the current density I to the reductant and oxidant concentrations at the electrode surface and to the driving force F(E − E 0).  F ( E − E 0 )   I  = k ( E ) [ Red]0 − [Ox]0 exp  −  F RT    

(13.2)

k(E) has then to be derived in the case of a CPET elementary step. Electrochemical CPET has been investigated theoretically by several groups [12–16]. The main item of CPET theories is a double Born–Oppenheimer approximation, which treats the electron as a fast subsystem with respect to the proton and treats the proton as a fast subsystem with respect to the degrees of freedom of the medium, as in proton transfer theories [17–21]. The four diabatic states are then mixed to generate two states that are adiabatic toward proton transfer as represented in Figure 13.2. In a CPET step, both electron and proton are transferred at the transition state corresponding to the crossing of these generated states where reactants and products have the same configuration. This configuration is reached through harmonic vibration of an environment bath representing the medium and describing the long-range electrostatic interaction of the system with a polarizable continuum and harmonic vibration of local dispersion modes, typically, the proton donor–acceptor vibration (Q mode) coupled to additional internal vibrations. It follows that the medium and the local mode interactions contribute to the reaction rate independently. The medium is treated classically and appears as a reorganization energy noted λ 0 in the expression of the reaction activation energy. At high temperatures with respect to internal vibration modes, their contribution also appears as a reorganization energy, noted λi, in the activation energy. Thanks to the harmonic approximation, the expression of the CPET rate constants thus coincides with the classical Marcus–Hush [22,23] formula for a simple nonadiabatic ET:  −λ k ( E ) = Z exp   4 RT 

 F (E − E 0 )   1 − λ  

2

   

(13.3)

in which Z is a preexponential factor detailed in the following text and λ = λ0 + λi is the total reorganization energy. These rate constants must be averaged over the electron energy in the electrode as in the Marcus–Hush–Levich development [24].

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486

Organic Electrochemistry Potential energy XOH..B

XR ..HB+

e–

XO ..HB+

XRH..B ZPE≠

+

H

H+ coordinate (q)

ΔG≠ E – E0CPET

XR ..HB+ XRH ..B

XOH...B

ZPER

XO ..HB+

H+ coordinate (q)

ZPEP H+ coordinate (q)

Heavy-atom reaction Coordinate

FIgURE 13.2 CPET pathway. Both reactant and product electronic states potential energies as function of heavy-atoms reaction coordinate are described by parabolas. Inserts show potential energies as function of proton coordinate.

The preexponential factor contains information regarding the coupling of the two electronic states, characterized by a coupling constant C. If C is small, a fully nonadiabatic regime is reached, and C = HET〈χi|χ f〉 where HET is the electron coupling constant and 〈χi|χ f〉 the overlap between the initial and final proton vibrational wave functions. A partially adiabatic transition takes place when the electron coupling constant HET is sufficiently large, whereas the resonance splitting of the proton levels remains small. The rate constant remains the same, but the coupling constant C is now described by a tunneling probability for the proton through a potential barrier: q   f  2π   2m p (V (q) − E ) dq   C (Q) = hν exp −  h     qi   ≠ 0

∫

where Q is the distance between the donor and acceptor atoms q is the proton coordinate ν 0≠ is the proton well frequency mp is the proton mass qi and q f are the classical turning points in each well at fixed Q

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487

Proton-Coupled Electron Transfers

It follows that the transition probability χ is a function of Q to be averaged according to χ=

∫

+∞

χ(Q)P(Q)dQ over the Boltzmann distribution P(Q). For a fully adiabatic transfer, the trans-

−∞

mission coefficient is 1. The link between these limiting cases is given by the Landau–Zener transition probability [25,26] χ, being given by χ=

2p 1+ p

(13.4)

In Equation 13.4, p is the probability of proton tunneling and ET taking place at the transition state as sketched in the upper insert of Figure 13.2. p is obtained from the Landau–Zener expression [27]:   C  p = 1 − exp  −π    RT  

2

πRT λ

   

(13.5)

Note that in a more refined development, electrochemical CPET rate constant expressions are derived that interpolate between nonadiabatic limits being defined in terms of weak vibronic coupling and fast solvent relaxation and solvent-controlled regimes defined in terms of strong vibronic coupling and slow solvent relaxation. In this chapter, we ignore solvent relaxation effects. As a result of the aforementioned derivation and (1) assuming that the electrochemical reaction takes place at a given distance from the electrode, (2) taking into account the multiplicity of the electrons’ electronic states in the electrode, and (3) considering the fact that the potential excursion in cyclic voltammetry does not exceed a few hundred millivolts thus allowing linearization of the quadratic term, the rate law (Equation 13.3) may be recast as [28]

k ( E ) = χk∞het

4πλ  −λ exp  RT  4 RT

 F (E − E 0 )   F (E − E 0 )   het exp k exp = S,CPET        2 RT   2 RT 

(13.6)

RT π RT het , times a factor accountk∞het is equaled to the collision frequency, kcoll = 2πM 1 + πRT /λ 4πλ ing for the multiplicity of electronic states in the electrode [24]. π  −λ   −λ  het het kShet exp  exp  = Z CPET ,CPET = χkcoll   is the standard rate constant, that 1 + πRT /λ  4 RT   4 RT  is, the rate constant at zero driving force (for the sake of simplicity, double-layer correction is not introduced here; it will be discussed later on). We thus end up with a rate law having the same formulation as for simple ETs. Besides the transfer coefficient α = 0.5, it is charac4πλ het = χk∞het the terized by two intrinsic parameters, λ the reorganization energy and Z CPET RT preexponential factor. The latter contains the kinetic characteristics of CPET relative to the proton transfer taking place at the transition state. These characteristics are discussed in details and illustrated by experimental examples in Section IV. Moreover, CPET pathways may be endowed with an H/D kinetic isotope effect (KIE) because tunneling of the proton is easier than tunneling of deuteron. However, the KIE is not an unambiguous diagnostic criterion in mechanism discrimination as shown in the next sections.

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III.

Organic Electrochemistry

COMPETITION bETWEEN STEPWISE (EPT AND PET) AND CONCERTED (CPET) PATHWAyS

The competition between stepwise and concerted pathway is a kinetics issue because all three pathways (EPT, PET, and CPET) have the same global thermodynamics characterized by EX0 O + BH+ /XR H + B or Eap0 . The competition can be easily analyzed for reactions taking place in buffered media when the electronic steps are rate determining and the proton transfers are unconditionally at equilibrium, which is often the case. Then the total current density I can be related to the electrode potential E, the apparent standard potential Eap0 , and an apparent standard rate constant kSap according to the following equation [29]:

(

 F E − E ap0 I ap = kS exp  F  2 RT 

)   Red 

(

 F E − E ap0  0 −  Ox  0 exp  − RT  

  

)    

(13.7)

The “apparent” character refers to the fact that Equation 13.7 does not represent an elementary step but it is a combination of standard rate constants characterizing each intervening pathway and also that it does not correspond to a rate constant at zero driving force but to the rate constant when the electrode potential is equal to Eap0 . Using the symbols defined in Scheme 13.4, the apparent standard rate constant may thus be expressed (Equation 13.8) as a weighted sum of the various standard rate constants, kSX, kSXH, and kSCPET − B, characterizing the PET, EPT, and CPET with B as proton acceptor pathways, respectively, [29]

kSap =

kSX [H + ] [H + ] 1+ 1+ K XO H K XR H

+

+

kSXH K O K R 1 + X+ H 1 + X+H [H ] [H ]

+

kSCPET − B K asO K asR [HB][B] K R [H + ] 1+ 1 + X+H K XO H [H ]

(13.8)

where HB/B (charge not shown) can be any acid–base couple present in the media. It thus appears that the apparent standard rate constant depends on the acidity constants, K XOH and K XR H, of the two acid–base couples, XRH/XR for the reduced species, and XOH/XO for the oxidized species. Figure 13.3 shows in a typical case the contributions of the various pathways to the apparent standard rate constant as a function of the pKs (considering, for the sake of simplicity, that only one acid–base couple participates to the concerted mechanism). It is seen that the smaller the pK XOH and the larger the pK XR H , that is, the higher in energy the reaction intermediates of the sequential routes, E 0XOH/XRH, k SXH XRH + B

K Ras

KXRH, KB

XOH + B

XRH ... B

E 0XO,H+ / XRH

k CPET–B S KXOH, KB XO ... HB O+ Kas

XR + HB

XO + HB

E 0XO/XR, k XS

SCHEME 13.4

PCET with equilibrated proton transfer.

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489

Proton-Coupled Electron Transfers 1

1

log(kSap)

1

log(kSap)

0

0

0

–1

–1

–1

–2

–2

–2

–3

–3

–3

–4

–4

pH

–4

pH

–5

–5 –4

0

4

8

12

0

4

8

12

16

(b)

(a)

pH

–5 –4

16

log(kSap)

–4

0

4

8

12

16

(c)

FIgURE 13.3 Variation of the apparent standard rate constant with pH as a function of the pK gap between redox states (a) pK XRH   = 9, pK XOH   = 2; (b) pK XRH   = 14, pK XOH   = −3; (c) pK XRH   = 16, pK XOH = −5). Black line: apparent standard rate constant; light gray: stepwise pathways contribution; dotted line: concerted pathway contribution. [B] + [HB] = 0.1 M; pKHB = 5. kSX = kSXH = 1 cm s−1; kSCPET − B K asO+ K asR = 0.01 cm s−1 M −1.

the more dominant the concerted (CPET) contribution. As can also be read both from Equation 13.8 and Scheme 13.4, the buffer plays a crucial role in the competition between stepwise and concerted pathways. It appears from Equation 13.8 that the more concentrated the acid–base couple, the stronger the contribution of the CPET (as illustrated in Figure 13.4). This effect is simply the result of the base being a reactant in the CPET process, whereas the acid–base couple serves only as a rapidly equilibrated pH buffer in the stepwise pathways. How do we know in practice whether one or the other mechanism takes place? Three main criteria may be used in this purpose. The first criterion derives from the variation of the apparent standard rate constant with pH and its adherence to Equation 13.8, provided that intrinsic kinetic and thermodynamic parameters are known or could be reasonably bracketed. Successful fitting using only contribution from stepwise pathways would be a strong indication that the reaction proceeds in two steps. A second clue is the dependence of the apparent standard rate constant from the acid– base couple concentration indicating the occurrence of a CPET pathway. A significant KIE would be a further indication of a concerted mechanism since it is the only one to involve proton transfer in the rate-determining step. It should, however, be mentioned that careful correction of possible 1

1

1

log(kSap)

log(kSap)

log(kSap)

0

0

0

–1

–1

–1

–2

–2

–2

pH

pH

–3 –4 (a)

pH

–3 0

4

8

12

–3 –4

16 (b)

0

4

8

12

–4

16

0

4

8

12

16

(c)

FIgURE 13.4 Variation of the apparent standard rate constant with pH as a function of the buffer concentration (a) [B] + [HB] = 0.5 M; (b) [B] + [HB] = 2 M; (c) [B] + [HB] = 5 M). Black line: apparent standard rate constant; light gray: stepwise pathways contribution; dotted line: concerted pathway contribution. pK HB = 5. kSX = kSXH = 1 cm s−1; kSCPET − B K asO+ K asR = 0.01 cm s−1 M −1; pK XRH = 9, pK XOH = 2.

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490

Organic Electrochemistry

thermodynamic isotope effects (differences in pKs for H and D) should be carried out before reaching a reliable conclusion [30]. In addition, proton transfers may not always be at equilibrium, as assumed earlier, especially in nonaqueous media opening the possibility of a KIE for the stepwise pathways. In this connection, it has been established that proton transfer can be considered as equilibrium if [31]  RT kdif  B  pK XR H − pK XO H < 0.5 + log    F v  

(13.9)

where v is the scan rate kdif is the diffusion limit rate constant [B] is the concentration of base The oxidative electrochemistry of a [OsII(bpy)2pyH2O]2+ complex provides a good example of a competition between stepwise and concerted pathways [29,32]. Two successive waves are observed in cyclic voltammetry corresponding to the passage from OsII(OH2) to OsIII(OH) and then to OsIV(O). For the OsII(OH2)/OsIII(OH) couple, the difference between the two pK involved is too small (pK XRH   = 9.1 and pK XOH = 2) for the mechanism to be stepwise, thus confirming that a small pK gap involves energetically inexpensive reaction intermediates and contains a negligible concerted contribution. Very large amounts of the buffer have to be added for the concerted pathway to start interfering (Figure 13.5). The acidic constants of the OsIII(OH)/OsIV(O) couple are not accessible in the experimental pH range, and the pK gap can thus to be assumed being higher than 14 [33]. In this case, the reaction follows a concerted (CPET) pathway as demonstrated by the variation of the apparent rate constant with pH, the variation of the apparent rate constant with the buffer concentration, and the observation of a significant KIE (Figure 13.6). This illustrates that the concerted mechanism is favored when intermediates of the stepwise pathways are high in energy. Oxidation of phenol in buffered water also illustrates competition between mechanisms but for a system that does not exhibit a reversible cyclic voltammetry wave [34], that is, a system in which

10

2.4

i (μA)

log kSap (cm/s)

1.6

–1

ap

log k (cm/s) S

5

–1.5

0.8

log[CH3COO–]

–2

0

0

–3

–2

–1

0

1

–0.8 –5 –1.6 E (V vs. NHE)

pH –2.4

–10 –0.2

0

(a)

0.2

0.4

0.6

0

2

4

6

8

10

12

14

(b)

FIgURE 13.5 Cyclic voltammetry of [OsII (bpy)2 pyH 2 O]2+ in a 0.1 M Britton–Robinson buffer. (a) Typical two-wave voltammogram at pH = 3. (b) Variation of the apparent standard rate constant with pH for OsII(OH2)/ OsIII(OH) couple. Circle: in H2O, square: in D2O. Insert: dependence of the apparent standard rate constant on buffer concentration in an aqueous acetate buffer at pH 5.

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491

Proton-Coupled Electron Transfers –2.5

log kSap (cm/s)

–3

–3.5

–4

–4.5

pH –5 3

4

5

6

FIgURE 13.6 Oxidation of [OsII (bpy)2 pyOH]2+ in a 0.1 M Britton–Robinson buffer. Variation of the apparent standard rate constant with pH. Circle: in H2O, stars: in D2O. Dotted line: prediction for a stepwise mechanism. Full line: prediction for a CPET pathway.

the species resulting from the e− + H+ transfer is not stable, being engaged in further reaction, viz., dimerization in the present case. At low scan rates, the electrochemical oxidation of phenol involves a fast and reversible proton-coupled ET followed, whatever its mechanism, by a rate-determining dimerization. Knowing the latter rate constant, its effect on the cyclic voltammetric responses can be corrected for so as to establish the Pourbaix diagram as shown in Figure 13.7. Assignment of the 0

1.8

Eap

EPET

0

1.6

ArOH +

0,H O

1.4

2 ECPET

ArO

1.2 1 0 EPET

ArOH 0.8 ArO–

0.6

pH 0.4 –10

–6

–2

pKArOH +

FIgURE 13.7 Pourbaix diagram for phenol in water.

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2

0

6

10

pKArOH

14

492

Organic Electrochemistry

PCET mechanism and characteristic rate constants can be achieved upon raising the scan rate. An apparent standard rate constant, kSap , is obtained from the variations of the peak potential with scan rate and phosphate concentration. kSap is a measure, according to Equation 13.8, of the superposition of all pathways, stepwise and concerted, involving B = H2O, HO −, HPO42−. Equation 13.8 may be recast, noting that the experiments are carried out at pH = 7.2, the pKa of hydrogen phosphate so as to make appear a 2− term independent from phosphate concentration, kSindep HPO4 , and a term proportional to phosphate concentration: kSCPET − HPO4

2−

kSap = kSindep HPO4 +

1 + 10

( pK

PhOHi +

− pH )

2−

× 1 + 10

( pH − pK PhOH )

HPO 4 2−   

(13.10)

The very fact that the apparent standard rate constant is a unity slope linear function of phosphate concentration points to the occurrence of a CPET-HPO42− pathway (Figure 13.8a). The relative contributions of the various pathways besides the contribution of the CPET-HPO42− pathway can be obtained by the application of Equation 13.8 and leading to the diagrams in Figure 13.8b. The EPT pathway contribution is very low because PhOH∙+ is high in energy (pK = −2). The PET contribution is the most important at pH > 9 as expected from the phenol pK (= 10) but the CPET-HPO42− pathway is dominating for pHs around hydrogen phosphate pK provided buffer concentration is high enough. The competition between stepwise and concerted mechanism can also be examined in nonbuffered medium, taking into account the diffusion of proton and OH− ions may interfere in the kinetics. Again, oxidation of phenol in unbuffered water can be used as a typical example [35,36]. The wave in basic pHs corresponds to the oxidation of phenoxide ion. It decreases with pH, as predicted by the PET process controlled by the diffusion of OH−, at the expense of a more positive wave, which is under partial control of the diffusion of the protons generated by the oxidation of phenol.

1.5

1

kSap (cm/s)

log kSap (cm/s)

0 –1 –2

1

–3 –4 –5 0.5

–6 –7 Phosphate buffer conc. (M)

–8

0

pH

–9 0

0.2

(a)

0.4

0.6

0.8

1

4

5

6

7

8

9

10 11 12

(b)

FIgURE 13.8 Electrochemical oxidation of phenol. (a) Apparent standard rate constant as a function of phosphate concentration at pH 7.2; (b) Contribution of the various pathways to the apparent standard rate constant. Dashed line: EPT; dashed–dotted line: PET; dotted line: CPET-H2O; dashed–double dotted line: CPET-OH−; full gray line: CPET-0.25 M HPO43−; full black line: CPET-0.5 M HPO43−.

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493

Proton-Coupled Electron Transfers

This wave corresponds to a CPET pathway rather than to an EPT pathway, which would require going to higher oxidation potential to reach the unstable cation radical intermediate. In nonaqueous medium, the CPET pathway is facilitated by the presence of a proton-accepting (or donating) group forming an H-bonded structure with the substrate being oxidized or reduced. Oxidation of phenols with attached proton acceptor is a typical example in which the CPET pathway prevails [37]. As mentioned in the introduction, such systems mimic a PCET process taking place in PSII in the charge transport pathway between the chromophore and the OEC. Using values bracketing the standard potentials and equilibrium constants of each step of the stepwise pathways, simulation of the cyclic voltammetry responses shows that stepwise pathways can be discarded. The reason is again that the intermediates are too high in energy. The observation of a significant KIE at scan rate high enough for the charge transfer to kinetically interfere in the process confirms the concerted character of the process. Formation of an H-bonded structure in the CPET process is not restricted to nonaqueous medium as illustrated by the OsIIIOH to OsIV=O CPET oxidation in buffered water [33]. The presence of an anion such as NO3−, able to bind with OsIIIOH, leads to the inactivation of the CPET pathway by preventing the formation of an H-bonded structure between OsIIIOH and the buffer base component. Conversely, formation of an H-bonded structure is not a sufficient condition for the CPET pathway to prevail (Scheme 13.5). Oxidation of guanine (GH) in aprotic solvent in the presence of cytosine (C) compared to its oxidation in the presence of 2,6-lutidine (L) gives an illustrative example of this point [38]. No H-bonded structure is formed between GH and L as checked by NMR. The oxidation pathway follows in this case an EPT pathway kinetically controlled by the initial ET step (∂Ep/∂logv ≃ 60 mV/decade). In the presence of C, GH forms a H-bonded (GH…C) structure (association constant of 600 M−1 measured by NMR) but, as with L, the electrochemical reaction is kinetically controlled by the initial formation of the cation radical as attested by the absence of displacement of the wave upon addition of C, the shape of the wave, its displacement with scan rate, and absence of KIE. This observation implies that the H-bond pairing does not favor the concerted pathway in this case. GH

C O

N N R΄

RO

NH2 NH

N

N NH2

N R˝ RO

O H H

H

O

O

H3C

O

O H

CH3

H

H

H

OR

H

R = tert-butyldimethylsilyl

H

R˝

R΄ GH L

O

N

C H2N

N N R΄

SCHEME 13.5

H-bonded structure.

© 2016 by Taylor & Francis Group, LLC

NH

N

NH2

O

N

N R˝

494

Organic Electrochemistry

IV. INTRINSIC CHARACTERISTICS OF CPET PROCESSES A.

SIMILARITY AND DIFFERENCES BETWEEN ET, CPET, AND CDET

As detailed in Section II, the rate law for an electrochemical CPET has the same form as rate laws derived for a simple outer sphere ET or concerted dissociative electron transfer (CDET, see Chapter 14) in which a bond between two heavy atoms is broken. It is characterized by two intrinsic het parameters, λ the reorganization energy and Z CPET the preexponential factor. Besides this general similar formulation, there are, however, important differences in these intrinsic parameters as detailed in the succeeding text. In all three cases, the reorganization energy is the sum of two contributions, a solvent reorganization energy λ 0 and an internal reorganization energy λ i. The internal reorganization is much larger in the case of CDET as compared to the two other cases because it includes the homolytic bond dissociation energy D (of the order of 2–3 eV) of the bond being broken as derived from the Morse curve model (see Chapter 14). This term is responsible for the kinetic penalty of the concerted dissociative ET pathway in competition with a stepwise pathway. The remaining internal reorganization due to all other interatomic distances and angle changes involving heavy atoms may be estimated quantum mechanically by calculating the energy of the product system in the configuration of the reactant system or vice versa. It is usually in the order of 0.5–1 eV. Regarding the solvent reorganization, its formulation based on electrostatic models is slightly different in the case of a CPET as in the two other cases [14]. Indeed, the solvent reaction coordinate for a CPET pathway is made of two ingredients, a fictitious charge number representing solvent reorganization upon ET, as in the case of an ET or a CDET process, and a dipole variation index representing solvent reorganization upon proton transfer. An electrostatic model sketched in Figure 13.9 allows the delineation of two independent contributions to the solvent reorganization energy noted λ 0ET (Equation 13.11) and λ 0PT (Equation 13.12). The first is identical to the solvent reorganization energy obtained for ET and CDET and is typically in the order of 1 eV. The second refers to the change in dipole moments accompanying CPET. It is usually small of the order of few tens of eV: λ 0ET =

1 1 e2  1 −   4πε0  εop εS  2a

z

z zR

zP

a HB

θ

B–

y

φ

a

A

y

φ x

e–

FIgURE 13.9 Modeling solvent reorganization in CPET.

© 2016 by Taylor & Francis Group, LLC

θ

μP

μR

x

(13.11)

AH

495

Proton-Coupled Electron Transfers

λ 0PT =

1 4πε0

 εS − 1   εop − 1   (µ R − µ P )2   − a3  2εS + 1   2εop + 1  

(13.12)

in which ε0 is the vacuum permeability εop and εS are the optical and static dielectric constants of the solvent a is the radius of the reactant equivalent sphere μR and μP are the dipole moments of the reactant and product, respectively Besides the reorganization energy, the second intrinsic kinetic parameter is the preexponential factor 4πλ het π RT het het Z CPET = χk∞het . k∞ is equaled to the collision frequency, kcoll , times a factor RT 1 + πRT /λ 4πλ accounting for the multiplicity of the electron electronic states in the electrode. This simplified model assumes that the electrochemical reaction takes place at a given distance from the electrode. A more refined treatment is available and will be described later on, but this simplified version can be used to compare ET, CDET, and CPET processes. The specificity of each process is described by the factor χ related to the probability p (Equation 13.4) for the system to jump from one electronic state to the other, which is obtained from the Landau–Zener expression depending on the coupling constant C between both states (Equation 13.5). For both ET and CDET, the coupling constant can be assumed to be large enough for the probability p and hence the factor χ to be equal to unity. The situation is different for a CPET process. Proton transfer occurs at the transition state between two vibrational proton states as sketched in Figure 13.2. The contribution of proton excited states may be additionally involved especially when the driving force of the reaction is large. Because the proton is much heavier than an electron, the coupling constant between the reactant and product states may be smaller in the case of a CPET as compared to ET and CDET. This is the kinetic penalty to pay in order to get the proton transfer concerted with ET and thus get benefits of the full driving force of the reaction. This a priori comparison of ET, CDET, and CPET properties indicates that (1) CDET has a larger reorganization energy than ET and CPET and (2) CPET has a smaller preexponential factor than ET and CDET. The intrinsic properties of CPET processes embodied in the preexponential factor can be analyzed by looking at the coupling constant. The coupling constant may be estimated from quantum calculation using the approximation that C = HET〈χi|χ f〉 where HET is the electron coupling constant and 〈χi|χ f〉 the overlap between the initial and final proton vibrational wave functions, or it can be estimated semiclassically using the tunneling probability for the proton through a potential barrier. A simple model, sketched in Figure 13.10, has been proposed that allows an estimation of Ceq (i.e., the coupling constant corresponding to the equilibrium distance between the proton donor and the acceptor) as a function of the barrier height ΔV depending on the distance Q between the proton donor and acceptor atoms [39]:  8 2 C (Q) = hν exp  −  3  ≠ 0

hν 0≠  ∆V ≠ 1  −   ∆V ≠  hν 0≠ 2 

3/ 2

   

with ∆V ≠ (Q) =

0 0  f0≠  Q − dAH − dDH   4  2 

where f0≠ = 4π2ν 0≠ 2 mp is the force constant of the proton well 0 0 dDH and dAH are the proton equilibrium distances in the reactant and product, respectively

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496

Organic Electrochemistry Potential energy

e– e–

ΔV≠ H+

H+

Proton coordinate

FIgURE 13.10 Modeling the proton tunneling barrier in CPET.

The actual coupling is thus a function of Q, so proton tunneling between the reactant and product states is a function of the donor–acceptor vibration with the shorter distance yielding easier proton tunneling. In agreement with the aforementioned model, the coupling constant can be described by the expression C (Q) = Ceq exp  −β(Q − Qeq )  in which the parameter β is the attenuation factor of the exponential decay of the vibronic coupling of the two states with Qeq and Ceq being the proton donor–acceptor equilibrium distance and the equilibrium coupling constant at equilibrium distance, respectively. In line with the proton being heavier than an electron, β values are in the order of 20 Å−1, whereas corresponding values are typically 1 Å−1 for ET. In a classical description, the contribution of each distance Q to proton tunneling is obtained by weighting the transmission coefficient by the Boltzmann probability P(Q) that the proton donor and acceptor atoms are at a distance Q from one another. In the context of a nonadiabatic CPET, this averaging leads to the following expression of the preexponential factor [40]:  2 RTβ2  het het = Z CPET Z CPET ,eq exp    f  which involves the combination of two intrinsic parameters: an equilibrium preexponential factor het Z CPET ,eq characterizing the coupling of electronic states in the transition state at the equilibrium proton donor and acceptor distance and a distance-sensitivity parameter β2/f in which f is the force constant of the harmonic oscillator of the proton donor and acceptor H-bond vibration. Full characterization of intrinsic properties of CPET processes thus requires determination of three parameters: the reorganization energy λ, the distance-sensitivity parameter β2/f, and the equilibrium het preexponential factor Z CPET ,eq . These three parameters may not be easily obtained separately from electrochemical experiments for various reasons, but combined approach of several techniques including electrochemistry [36], photoinduced ET [41], and stopped-flow [42] allows, as shown in the following example of phenol oxidation in water, to fully characterize a CPET process and check the consistency of both heterogeneous and homogeneous CPET kinetics.

B.

HYDROGEN-BONDED SYSTEMS

To take advantage of their favorable thermodynamics, CPET pathways need not be too severely penalized kinetically in order to prevail over the competing stepwise pathways. Efficient proton

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497

Proton-Coupled Electron Transfers R2

H

H N

O

O

tBu

H N

O

tBu R2 R1

tBu

tBu

AP

SCHEME 13.6

R1

R2

1

CF3

CH3

2

H

CH3

3

CH3

CH3

4

H

H

H-bonded amino-phenol systems (AP).

tunneling is thus required as described earlier, which implies short distances between the group generating the proton upon oxidation and the proton acceptor (and vice versa for a reduction process)— a necessary albeit not a sufficient condition as shown in the case of guanine oxidation in aprotic solvent. Molecules containing an oxidizable phenol moiety and an attached nitrogen base serving as proton acceptor have been intensively investigated to get insights into the process. Measurement of the standard rate constant at various temperatures for the aminophenol noted AP (Scheme 13.6) allows getting both the preexponential factor and the reorganization energy from an Arrhenius plot (provided double-layer effect and zero-point energies are being taken into due account) [43]. The reorganization energy is around 1.4 eV in agreement with an estimation of the solvent reorganization energy of about 1 eV and a calculation of the intramolecular reorganization energy of het 0.4 eV. The preexponential factor is 34,580 cm s−1, much larger than expected from the factor kcoll thus indicating that a realistic analysis has to take into account the fact that the reaction may take place at various distances from the electrode surface similarly to what happens with simple ET. A complete analysis of the problem taking into account likely approximations leads to the conclusion that the rate law remains unchanged as well as the expression of the standard rate constant. The rate constant appears as the product of a preexponential factor and an exponential factor depending on the reorganization energy [43]. However, the preexponential factor for proton donor and acceptor at equilibrium distance is now het Z CPET ,eq =

νn βe

2   2π2Ceq π ,0 ln  1 +  πRT  hν n 4πRTλ  1+ λ

where νn is the frequency vibration of the activated complex along the reaction coordinate (typically 1012 s−1) βe is the decay constant for the coupling constant with the distance between the electrode and the reactant (typically 1 Å−1) [44] Ceq,0 is the CPET coupling constant for proton donor and acceptor at equilibrium distance and the reactant system at minimal approach distance from the electrode

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498

Organic Electrochemistry

In the nonadiabatic limit (corresponding to the case of interest), the averaging over Q distances leads to het Z CPET =

1 βe

2 π2 π πRT h 4πRTλ 1+ λ

 2  4 RTβ2   Ceq,0 exp    f   

het from the intercept Knowing the reorganization energy from the Arrhenius plot slope and Z CPET 2  2 RTβ  allows to get an effective coupling constant: C0,eff = Ceq,0 exp   = 0.015 eV (0.007 eV for the  f  deuterated AP compound indicating a H/D KIEs in line with the concerted process). Independent assignment of the two intrinsic parameters Ceq,0 and β2/f could not be achieved. However, comparison with a simple outer sphere ET, whose preexponential factor is 54,000 cm s−1 (corresponding to an outer sphere ET coupling constant of 0.018 eV), indicates that the proton tunneling is efficient for this heterogeneous CPET process. It is interesting to note that a similar treatment of homogeneous results obtained with a similar aminophenol molecule leads to the conclusion that C0,eff is smaller in the homogeneous case. This is deemed to derive from the effect of the strong electric field within which the electrochemical reaction takes place. Nonetheless, the tight H-bonded structure is crucial for the efficiency of the CPET process. Homogeneous studies on similar aminophenol systems have shown a strong rate dependence on increased proton transfer distance [45]. Computational results indicate that anharmonicity of the H-bond vibration and influence of proton vibrational excited states may have to be considered to describe these systems, thus illustrating the complexity of CPET processes [46]. Despite these complications, it remains that the distances over which the proton may travel as a result of a CPET reaction appear to be limited to the rather small values of H-bond length (ca. 2–3 Å). However, the idea according to which this distance might be substantially increased by inserting an H-bond relay between the group being oxidized (or reduced) and the distant proton acceptor (or donor) has been explored. A series of molecules containing an oxidizable phenol and a pyridine group that serves as proton acceptor and an alcohol function between them (Scheme 13.6) have allowed testing experimentally the concept of H-bond relay [47,48]. In all cases, a chemically reversible wave is obtained. Thermodynamics arguments can be used to discard any stepwise mechanism, and it is concluded that the displacement of the two protons is concerted with ET in line with the observation of a KIE. It is worth noting that x-ray data show that the structure is not folded and that the distance between the proton donor and acceptor sites is ca. 7 Å. Similar analysis of cyclic voltammetry data than that performed with AP indicates that the reorganization energy is almost constant over the whole series of compounds while the standard rate constants of the four H-bond relay molecules are much smaller than the standard rate constant of AP (Table 13.1). Thus, as expected and confirmed by computational studies, the reason that makes CPET oxidation of the H-bond relay molecules intrinsically slower than the oxidation of AP is that the tunneling

TAbLE 13.1 Reorganization Energies and Standard Rate Constants for Molecules in Scheme 13.6; see Text for the Definition of Parameters Molecule

λi (eV)

kS (cm s−1)

AP 1 2 3 4

0.390 0.405 0.433 0.420 0.443

8 × 10−3 9 × 10−4 4.5 × 10−4 8 × 10−5 4.5 × 10−4

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499

Proton-Coupled Electron Transfers e–

Potential energy

H+

XRH1 ..RH2 ..B ZPE≠

XO ..H1R ..H2B qH2 qH1

ΔG≠ F E – E0

XR H1 .RH2 ...B

qH2

ZPER qH1

XO ...H1R .H2B qH2

ZPEP

qH1 Heavy-atom reaction coordinate

FIgURE 13.11

Modeling CPET in H-bonded systems with an H-bond relay.

efficiency is less in the first case because two protons move concertedly with the electron as sketched in Figure 13.11. It is thus demonstrated that a net displacement of a proton in concert with an ET over a distance as large as 4.5 Å is achieved thanks to an H-bonded relay. The kinetic penalty for this is that the coupling constant and hence the preexponential factor are decreased [49]. The decrease by a factor of ca. 50 of the standard rate constant when the traveling distance of the proton is increased by ca. 2 Å indicates that the distance-sensitivity factor β is of the order of 2 Å−1, which is large compared to corresponding values for simple proton transfer (ca. 20–30 Å−1). Modeling of such a 2D CPET will require further development including the influence of the H-bond-accepting and H-bond-donating properties of the relay on the potential energy surface. The observed variations of the standard rate constant within the series (Table 13.1) are indeed likely due to the modification of these properties by the substituent on the relay alcohol. The illustration of the possible translocation of proton over a long distance through H-bonded relays concertedly with ET indicates that similar process can be envisioned with proton transfer along a water chain, as exemplified by the reduction of superoxide ion with water as proton donor [50]. In an aprotic solvent, dioxygen exhibits a quasi-reversible cyclic voltammetry first wave followed by a second, broad irreversible wave. This second wave corresponds to a CPET reduction of the superoxide ion. Indeed, if a PET pathway was followed, the first wave should involve two electrons and, if an EPT pathway was followed, the second wave would be sharp (with a transfer coefficient lower than 0.5). The observation of a significant KIE (KIE = 2.5) also confirms the occurrence

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Organic Electrochemistry O

O– OH2 e–

O

O–

O–

H

H2 O

H 2O

H

O H

O H

e– O

O–

H

H O H

H

O

O–

O–

H

H

O H

H O H2O

O–

H 2O

H

H

O H

e– O–

O

O H

H

O–

O

H

H O

H H

H O

–O

H

H2O

H2O H

H

e–

O H

O H O

O–

O–

H

H

H

O H

H

H O

H

O

H O

SCHEME 13.7

H

O

O

H

H

O H

H –O

Reduction of superoxide ion with water as proton donor.

of a CPET pathway. The large anodic shift of the second wave upon addition of water [51] has been interpreted as resulting from a concerted transfer of one electron and one proton through short water chains (Scheme 13.7). The main factor driving the reaction toward such a kinetically demanding process is the thermodynamic advantage due to a decrease of the attending repulsion between HO2− and OH−. The same behavior is observed for the reduction of the benzophenone radical anion upon addition of water [52]. The ability of water to form H-bonded chains creates a new type of CPET pathway allowing the transport of a proton over a long distance concertedly with ET. This may be useful for a better understanding of biological reactions involving PCET such as in the disproportionation of superoxide by superoxide dismutase [53].

C. WATER (IN WATER) AS PROTON ACCEPTOR If water may serve as proton relay as demonstrated in Section IV.B, it is also an ubiquitous proton donor and acceptor. Its role in PCET reactions as a peculiar, H-bonded and H-bonding, proton donor and acceptor when it is used as solvent is not only an important fundamental issue but is also of considerable interest for the comprehension of natural systems. Although investigated over decades, the mechanisms of proton conduction in water are still under active experimental and theoretical scrutiny [54,55]. The question raised here is whether or not proton transport in water can be concerted with ET. In other words, does the peculiar character of water induce specific intrinsic properties of CPET with water (in water) as proton acceptor? Interesting insights have been obtained through the investigation of phenol oxidation using several techniques, electrochemistry and also laser flash photoinduced ET and stopped flow techniques.

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Proton-Coupled Electron Transfers 12 i (μA)

10 8 6 4 2 0

E (V vs. NHE) –2 0.5

0.6

0.7

0.8

0.9

1

1.1

1.2

1.3

FIgURE 13.12 Cyclic voltammetry of phenol in water at 0.2 V/s in unbuffered water at various pHs (from right to left: 2, 3, 4, 5, 6, 7, 8, 8.5, 9, 9.5, 10, 11, 12).

A first kinetic characterization of phenol oxidation with water as proton acceptor can be obtained from cyclic voltammetry in unbuffered water [35]. Once the pH is low enough so that the OH−-PET pathway is too slow, the followed pathway is a CPET with water as proton acceptor. If the initial proton concentration is large (pH < 3), then the wave is kinetically controlled by dimerization of the phenoxyl radical. At pH > 3, the initial concentration of protons is perturbed by their production from phenol oxidation (Figure 13.12). It is thus demonstrated that the wave depends on a single dimensionless parameter p′ measuring the competition between dimerization and CPET kinetics: p′ =

kSCPET − H2 O (C 0 /CS )1/ 2 ( DPhOH )1/ 4 ( DH+ )1/ 4 ( Fv /RT )1/ 3 (4kdimC 0 / 3)1/ 6

in which C0 is the phenol concentration CS is a normalizing concentration taken equal to 1 M Raising the scan rate allows determination of the standard rate constant kSCPET− H2 O = 25 cm s−1. The measurement of such a high standard rate constant at such moderate scan rates can be achieved because the follow-up dimerization is fast and competes with the “termolecular” backward CPET, PhO • + H+ + e−. Because the medium is unbuffered, protons produced by the forward CPET diffuse away from the electrode as attested by the appearance of the proton diffusion coefficient in p′ making the competition less in favor of an equilibrated CPET followed by a rate determining dimerization as observed at pH < 3. Repeating the experiments in deuterated water allows measuring a KIE of 2.5, in line with the concerted character of the reaction. Deciphering the reasons behind such a large standard rate constant requires separating the various intrinsic ingredients of kinetics, that is, the reorganization energy, the distance-sensitivity parameter β2/f, and the equilibrium preexponential factor. Homogeneous techniques have appeared to be suitable for that purpose [40]. Measurements of the CPET homogeneous rate constant for phenol oxidation as a function of the driving force by varying the electron acceptor standard potential leads to the reorganization energy (Figure 13.13). The major part of this reorganization energy may be ascribed to solvent reorganization accompanying the generation of a water-solvated proton. It is noteworthy that the ensuing value of the

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Organic Electrochemistry 12 log k 10 8 6 4 2 0 –2 ΔG0 (eV) –4 –1

–0.5

0

0.5

FIgURE 13.13 Forward rate constant of phenol oxidation by various electron acceptors. Black and gray symbols: in H2O and D2O, respectively. Squares: RuIII(bpy)(4,4′-CO2Et-bpy)2; triangles: RuIII(bpy)3; circles: RuIII(4,4′-methyl-bpy)3 ; tilted squares: IrIVCl6. Line: activation-controlled rate constant predicted for a reorganization energy of 0.45 eV. CPET,H 2 O self-exchange reorganization energy, λ se = 0.45 eV, is remarkably small. The proton charge is accordingly not concentrated on a single hydrogen atom or even on a single-protonated water molecule. It spreads over a ca. 6.5 Å-radius water cluster molecules, in agreement with recent spectroscopic observations [56] and with the aforementioned description of the ability of the proton to move over several H bonds in concert with ET. With the reorganization energy in hand, analysis of temperature dependence of rate constants gathered from laser flash experiments leads to the values of the distance-sensitivity parameter β2/f and the equilibrium preexponential factor in both light H2O and D2O. As expected, the equilibrium preexponential factor is smaller with deuterium than with hydrogen, and the distance-sensitivity parameter is larger with deuterium than with hydrogen (Table 13.2). These values are compatible with a Grotthus-type mechanism during the CPET process intrinsically efficient. Indeed, comparison with other proton-accepting bases in water, hydrogen phosphate, and pyridine [57] (Table 13.2) shows that the distance-sensitivity parameter is much smaller in the latter cases than with water due to a much stiffer phenol-based system and,

TAbLE 13.2 Kinetic Parameters for CPET in Water with Water, Phosphate, or Pyridine as Proton Acceptor; see Text for the Definition of Parameters H2 O

HPO42−

Pyridine

λCPET (eV)

0.45

0.86

0.53

β2 (K−1) 2R f

0.0125 (H) 0.020 (D)

0 (H) 0.0064 (D)

0 (H) 0.006 (D)

ln (Zeq)

19.5 (H) 15.9 (D)

16.8 (H) 14.8 (D)

17.6 (H) 15.1 (D)

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Proton-Coupled Electron Transfers

in this sense, a less efficient CPET because the proton does not travel over a large distance in concert with ET. The third intrinsic parameter, the equilibrium preexponential factor, is also larger in the case of water than in the case of hydrogen phosphate and pyridine, indicating that proton translocation is more efficient in water than it is for a conventional CPET process where the proton is more localized. These results are in qualitative agreement with the electrochemical data that showed that the phenol oxidation CPET standard rate constant corrected for double-layer effect,  (2 z + 1)FφS  −B kSCPET = kSCPET − B exp  ,corr  RT   (z is the charge number of the reactant system and ϕS the potential at the reaction site vs. the solution − H2 O bulk potential) is much larger when water is the proton acceptor (kSCPET = 83 cm s−1) as com,corr CPET− HPO 42− = 0.002 M −1 cm s−1). pared to the case where hydrogen phosphate is the proton acceptor (kS A more quantitative comparison between homogeneous and electrochemical kinetic characteristics can be achieved as follows [34]. Knowing reorganization energies for CPET self-exchange reactions, CPET , the corresponding reorganization energies for electrochemical heterogeneous reactions are λ se estimated through a CPET  = λ se λ CPET el 1 − d    where a is the radius of the equivalent sphere d = 2(δH2O + a) the distance between the reaction site and its electrical image in the electrode (Figure 13.14) Then, the electrochemical preexponential factor is obtained from λ  kS,corr = Z het exp  − 4 RT 

π λ    exp  −  = Z el  4 RT 1 + πRT /λ   

(13.13)

Application to the comparison of the electrochemical and homogeneous results for H2O vs. HPO42− as proton acceptor indicates that the ratio of preexponential factors is substantially larger, by a factor of ca. 10, in the electrochemical case than in the homogeneous case. This may be due to the existence of a strong electric field at the reaction site favoring the zwitterionic form of the reactant system (PhO −, H+, n H2O) in the transition state. It is finally worth noting that, as expected, preexponential factors Z el for CPET reactions are smaller than the comparable factor of a simple outer sphere ET. In the case of the CPET phenol oxidation with water as proton acceptor, comparison to phenolate oxidation shows a decrease as high as five orders of magnitude. This decrease cannot be fully understood with the present model in which the preexponential factor is a measure of activationless reactivity of a reactant system assimilated as a sphere; otherwise, very large KIEs would be expected. The reactant system is in fact likely to be structured so as to adopt a precise spatial conformation allowing the formation of one or several H bonds as required by the occurrence of the CPET reaction. Both electrochemical and homogeneous data are thus consistent and show that water (in water) is an intrinsically very efficient proton acceptor in CPET processes.

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Organic Electrochemistry

–

O

H O

N

H

O

O

H O H H

H HO

O

H H O

H OH

H H O H H H O H O O H O H H H OH H O H H H H O H H O HO O H H H H O H O H H O H O H H O H H H H H O H H OH O H H O H HO O H H H H O H O H O H O H +H H H O H O HO H H H H H O H O O H H H O H H HO

H

H

e–

H

H O

O

H H O H H O

O–

H H

N O

O

O

H H

O– O

N

H H

O

O

O

O –

H H

N O O

O

H H O

–O

H

H P O

O

H

e–

O

O–

H H

H O

O

H H O H

O–

H O

O H

P

O

H –

O H

H

Potential

O H H O H H O H H O O H

E

O

O – N O

φ2 = 125 φHPO

2– 4

= 108

φH O = 53 2

0 δH O = 1.5 OHP 2

aNO– = 2.9 3

aHPO2– = 3.5 4

aH O = 6.5 2

FIgURE 13.14

Sketch of the electrode vicinity for phenol oxidation in water.

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Proton-Coupled Electron Transfers

V. CPET FOR bOND ACTIVATION A.

BREAKING BONDS WITH PROTONS AND ELECTRONS

The concerted dissociative ET theory (see Chapter 14) together with the numerous illustrating experimental examples shows the way in which electrons can break heavy-atom bonds [6]. Associating proton transfer with ET might be additionally profitable to break bonds because of the thermodynamic advantage of concerted pathways over stepwise pathways. Indeed, various pathways are conceivable for going from reactants to products as sketched in Scheme 13.8 in the case of a reduction (transposition to oxidation is straightforward) [58]. There is a competition between a three-step stepwise pathway, two combinations of concerted reactions with an additional step and an all-concerted pathway. The latter is the best way to make use of the thermodynamic advantage provided this is not counterbalanced by a large kinetic penalty. Combining the models previously developed for concerted dissociative ET and concerted proton– electron transfer leads to a kinetic model for the all-concerted process. As in the simple CPET case, two successive applications of the Born–Oppenheimer approximation lead to define the transition state in terms of a heavy-atom coordinate and the preexponential factor in terms of a proton displacement coordinate. Concerning the first of these applications, heavy-atom reorganization involves the solvent molecules, the vibrations of reactant bonds not being cleaved in the reaction, and, most importantly, the contribution of bond cleavage. Regarding the latter, it seems appropriate to use the same approximation for the potential energy curves as for concerted dissociative ETs with no accompanying proton transfers, that is, a Morse curve for the reactants and a repulsive Morse curve for the products (equal to the repulsive part of the reactant Morse curve) as shown in Figure 13.15. Consequently, the rate law of the irreversible all-concerted reaction is ≠  ∆Ghet I 3 rd = Z het exp  −  RT F 

  ×  Y−X  × HB  

with ≠

∆Ghet

λ + D  F (E − E 0 )  = het 1 +  4 λ het + D  

2

Bond breaking – X + HB Y—

Concerted electron transfer and bond breaking

Outer sphere electron transfer

Y X– + HB

Proton transfer Concerted bond breaking and proton transfer

e–, ED

Y XH + B–

Y — X + HB All in concert e–, ED

SCHEME 13.8

Mechanistic pathways for dissociative electron proton transfers.

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Organic Electrochemistry

Potential energy Y–XH+ ..–B e– H+

Y X– ..HB Y–X .. HB

H+ coordinate (q)

Y XH ..–B

ΔG≠

ΔG0

Y–XH+ .. –B

Y X– ..HB

Y–X .. HB

Y XH ..–B

H+ coordinate (q)

H+ coordinate (q)

Heavy-atom reaction coordinate mostly the Y–X distance

FIgURE 13.15

Modeling concerted dissociative proton electron transfers.

[Y–X] and [HB] are the concentrations of the indicated species at the electrode surface. The reorganization energy, λhet, includes the energy for solvent reorganization and internal reorganization in the Y–X molecule, besides the cleavage of the bond, E0 is the standard potential relative to the all-concerted reaction and E the electrode potential. As in the case of simple CPET reactions, the preexponential factor combines the formation of the precursor complex, the degree of adiabaticity of ET, and the effect of proton tunneling at the transition state. As for simple ET, CDET, or CPET, all electronic states can be taken into account. Also, as in the case of a CPET, the third-order character of such reactions does not prevent their occurrence, as shown, for example, by the oxidation of phenol with hydrogen phosphate or pyridine as the proton acceptor. In systems where the proton donor is attached to the structure that bears the cleavable heavy-atom bond, the rate law is replaced: ≠  ∆Ghet I 2 nd = Z het × exp  −  RT F 

  ×  Y−X, HB  

An all-concerted mechanism is thus characterized by a large reorganization energy due to the contribution of D and a preexponential factor affected by proton tunneling at the transition state. As described for CPET reactions, this last feature may induce an H/D KIE. However, the large irreversibility caused by the breaking of the heavy-atom bond implies that, in a cyclic voltammetric experiment,

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Proton-Coupled Electron Transfers H O O

O O

SCHEME 13.9 H-bonded peroxide molecule.

the driving force at the peak potential is large to overcome this high intrinsic barrier. It follows that the transition state closely resembles the initial state. One consequence is that, in the heavy-atom transition state, the •Y–X− moiety is much less basic than in the products’ geometry because the Y–X bond is barely broken. The result is that the intersection of the proton energy profiles of reactants and product electronic states in the transition state is likely to be close to the zero-point energy level. This indicates that the overlap of proton vibronic states is large and therefore insensitive to isotope substitution. Then the H/D KIE is predicted to be negligible in spite of the concerted character of the reaction. Moreover, the preexponential factor is not expected to be much smaller than preexponential factors of corresponding dissociative ET not concerted with proton transfer.

B. ACTIVATION OF MOLECULES That concerting proton transfer to electron can be profitable to activate a molecule through bond breaking has been illustrated with the cleavage of an O–O bond helped by the presence of a proximal carboxylic acid group (see Scheme 13.9) [58]. The kinetic response to the increased driving force due to the concerted proton transfer is revealed by the fact that the cyclic voltammetric wave of the acid is located at a potential less negative by 700 mV compared to the methyl ester. That the cleavage is concerted with ET in both cases is attested by the large width of the waves indicating a small value of the transfer coefficient. Thus, if a two-step pathway was to be followed with a first irreversible concerted ET and bond-breaking step followed by a downhill protonation step, then the kinetics of the reaction would not respond to the increase of driving force offered by the follow-up protonation. The all-concerted mechanism is thus operating and leads to a considerable acceleration of the O–O bond activation.

VI. CONCLUDINg REMARKS Proton-coupled ETs are omnipresent in natural and artificial chemical processes. Understanding mechanisms and activation–driving force relationships, which underlie their practical efficiency, is a timely task in front of contemporary challenges concerning energy conversion. Focusing on PCET reactions in which, in contrast with hydrogen-atom transfers, proton and electron transfers involve different centers, the reaction may go through an electron or proton transfer intermediate, giving rise to an EPT and a PET pathway, respectively. CPET reactions, in which proton and electron transfers are concerted, have the advantage of bypassing the high-energy intermediates of the stepwise pathways, even though this thermodynamic benefit may have a kinetic cost. Kinetics-based mechanism analysis is now available, making it possible to distinguish the three pathways and to uncover the factors governing the competition between, thanks to the modeling of the concerted pathway. Many illustrating experimental examples have been described, most of them inspired by biological systems, for example, PSII and superoxide dismutase. It has been also shown that water is a remarkable proton acceptor endowed with very small reorganization energies. Coming from different corners of chemistry, these examples reveal general features of PCET reactions. Further development concerns the coupling of PCET reactions with bond-breaking/bond-forming processes likely to be involved in the catalytic systems currently designed to activate small molecules (O2, H+, H2O, CO2, RCl) as required by the resolution of contemporary energy challenges.

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14

Dissociative Electron Transfers Cyrille Costentin, Marc Robert, and Jean-Michel Savéant

CONTENTS I. Introduction ............................................................................................................................. 511 II. Modeling Concerted Dissociative Electron Transfers ............................................................ 513 A. Morse Curve Model ........................................................................................................ 513 B. Sticky Concerted Dissociative Electron Transfer ........................................................... 515 III. Competition between Concerted and Sequential Dissociative Electron Transfers................. 517 A. Structural, Electronic Factors and Solvent Effects Controlling the Competition ........... 517 B. Role of the Driving Force ............................................................................................... 521 C. Electrochemical versus Photoinduced Dissociative Electron Transfer ........................... 523 IV. Cleavage of Primary Radicals ................................................................................................ 524 V. Concluding Remarks............................................................................................................... 528 References ...................................................................................................................................... 529

I. INTRODUCTION The coupling between electron transfer and bond breaking between two heavy atoms occurs in a large number of biochemical and chemical processes, such as cleavage of C–halogen bonds in organic halides as well as other bonds [1–4], electron transfer activation of small molecules involved in contemporary energy challenges (e.g., O2 and CO2), as well as enzymatic reactions like dechlorination processes of RCl toxic derivatives within reductive dehalogenases [5]. In all of these reactions, the bond breaking accompanying electron transfer may be triggered in various ways, electrochemically, by homogeneous electron donors or acceptors, photochemically, or by means of pulse radiolysis [1–3]. The fact that so many chemical reactions can follow or accompany electron transfer is the basis of the synthetic value of electron transfer chemistry. Such processes also irrigate more applied fields, for example, the area of sensors and biosensors, which both involve the transduction of the presence of a molecule into an electrochemical signal. Another, more prospective field, concerns molecular electronics, where the understanding of the structural changes coupled to charge transfer will be central in the design and working of devices including redox centers connected by molecular wires. The key reactivity paradigm in charge transfer– induced bond breaking processes is the concerted/stepwise mechanistic dichotomy, as illustrated in Scheme 14.1 [1–3]. Charge transfer and bond cleavage reaction may indeed occur concertedly according to a single elementary step (concerted dissociative electron transfer [CDET]), or in two successive steps, the electron transfer then leading to a frangible species that cleaves in a distinct step (sequential dissociative electron transfer [SDET]), as shown in the potential energy profiles of Figure 14.1. As we define it here, the term dissociative charge transfer thus embraces all reactions for which an electron transfer triggers the breaking of bond between two heavy atoms, irrespective of the exact pathway, concerted or stepwise. This is not always the case in the literature where dissociative means concerted in a number of cases. The context will allow one to distinguish between a concerted dissociative reaction and a stepwise dissociative one. Concerted reactions (CDET) have the

511

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Organic Electrochemistry Stepwise RX + e–

RX

–

(π anion radical)

Concerted (R , X–)

(σ anion radical)

R + X– e–: Electrode, homogeneous donor (ground state, excited state)

SCHEME 14.1

Dissociative electron transfer mechanisms.

Potential energy RX –E 0 RX/RX –

–E

–

Stepwise

RX + e– Concerted

–E 0

RX/R + X–

–E

RX + e–

R , X–

R + X–

Reaction coordinate

FIgURE 14.1 Potential energy profiles for the concerted and stepwise pathways according to Scheme 14.1. The E 0s are the standard potentials of the subscript couples. E is the electrode potential or the standard potential of the homogeneous donor.

thermodynamic advantage over a sequential process (SDET) of avoiding the formation of highenergy intermediates (i.e., the anion radical in the case of the reduction of a neutral substrate as shown in Figure 14.1), thus going directly, during the rate-determining step, to the final products. This advantage may, however, be counterbalanced by the large intrinsic activation barrier deriving from bond cleavage that results in a kinetic penalty as compared to a two-step process. The CDET modeling is presented in Section II. After the cleavage has taken place, another energy minimum is presented in Scheme 14.1 and in the potential energy diagram of Figures 14.1 and 14.2, corresponding to an ion–radical adduct that may or may not survive in a polar solvent, for both CDET and SDET mechanisms. When existing, this adduct, resulting from a charge–dipole attractive interaction between the cage fragments before they diffuse out, may be viewed alternatively as a σ* anion radical or as forming a weak three-electron bond. If such interactions are expected to decrease or even to vanish in polar liquid, experimental studies have confirmed their existence, at least when a partial positive charge is induced on the remaining radical part, thanks to the presence of a strong electron-withdrawing substituent. The consequences of these interactions on the dissociative electron transfer kinetics are discussed in Section II. The competition that exists between the CDET and SDET pathways depends upon intramolecular (structural, electronic) and environmental (solvent, energy of the incoming electron) factors.

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Dissociative Electron Transfers Potential energy

R + X– (R , X–)

(a)

R + X– (R , X–)

Length of the cleaving bond (b)

FIgURE 14.2 Energy as function of the intramolecular reorganization for the reduced system (RX + 1e−), with (dotted line) and without (full line) interaction between the fragments, along a sequential with an energy •− minimum at short distances corresponding to a RX intermediate (left, a) and a fully concerted (right, b) pathway.

A detailed analysis of the thermodynamic and kinetic impacts of these factors and the outcome of the competition are presented in Section III. As illustrated in Scheme 14.1, dissociative electron transfers may be triggered photochemically. In these cases, the competition between CDET and SDET pathways hinges upon the same factors as those in electrochemical reactions, although the driving force offered to the reaction is larger. These reactions are presented in Section III also. Sequential cleavage of ion radicals may occur in a homolytic or heterolytic manner, and in both cases, the cleavage amounts to an intramolecular dissociative electron transfer reaction. These aspects will be detailed in Section IV, with emphasis put on the systems where symmetry restrictions lead to the formation of an avoided crossing. For all of these various aspects of charge transfer–bond breaking– induced processes, systematic analysis of reaction mechanisms as well as the factors that control the competition between the various reaction pathways leads to structure–reactivity relationships and therefore to predictive rules. For all of these reactions in which electrons are used to break bonds, a rationalized description emerges from the analyzed experimental examples and from the mechanistic and theoretical analyses. Dissociative electron transfer leading to the cleavage of a bond linking two heavy atoms is also very often coupled to a proton transfer, as it is the case, for example, during the reduction of O2 into H2O. In these cases also, the degree of concertedness between charge transfer, proton transfer, and bond cleavage is essential in determining the kinetics and the efficiency of the reaction. These processes will be briefly introduced in Section V and fully discussed in Chapter 13. They are indeed involved in the activation of small molecules (O2, H+, H2O, CO2) for the storage of electricity (originating, e.g., from solar energy) into chemical bonds, a key challenge for the twenty-first century chemistry.

II.

MODELINg CONCERTED DISSOCIATIVE ELECTRON TRANSFERS

A. MORSE CURVE MODEL Potential energy curves describing both reactant and products are modeled by Morse curves, with the assumption that the repulsive interaction of the two fragments formed upon charge transfer is identical to the repulsive part of the reactant Morse curve [6]. Attending solvent reorganization is estimated from the Marcus–Hush model. These two ingredients of the model

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lead to a quadratic activation (activation free energy: ΔG ≠)–driving force (minus standard free energy: −ΔG 0) relationship as given in the following: ∆G ≠ =

DRX + λ 0  ∆G 0  1 +  4 DRX + λ 0  

2

(14.1)

where D RX is the homolytic bond dissociation energy λ0 is the solvent reorganization energy The standard free energy of the reaction leading to complete dissociation (E: elec0 trode potential, ERX/R : standard potential of the RX/R• + X− couple) is given by • +X− 0 0 ∆G = F ( E − ERX/R• +X− ) = F ( E + DRX − T∆S 0 − EX0 • /X− ), where ΔS 0 is the bond dissociation entropy • and EX0 • /X− the standard potential of the X /X− redox couple. When necessary, additional sources of intramolecular reorganization may be included as an additive term to the intrinsic barrier ∆G0≠ = ( DRX + λ 0 ) / 4 . The homolytic bond dissociation energy represents the kinetic penalty for the concerted reaction as compared to the sequential pathway. The entropic term in the driving force for the reaction could be estimated either from known thermodynamical data or by quantum calculation and usually represents a minor contribution. The electron transfer rates may then be expressed as in the Marcus–Hush theory [7–11]:  0 F E − ERX/R D + λ 0  + X− 1+ k ( E ) = Z exp  − RX 4 RT  DRX + λ 0  

(

•

)   2

   

where Z is the pre-exponential factor. If one assumes that the electron transfer takes place at a fixed distance from the electrode and including the effect of the multiplicity of electronic states of the electrons in the electrode leads to

Z = Z coll =

( RT )2 8 M (1 + πRT /λ)

where M is the molar mass of the reactants. One important consequence of taking all electrode electronic states into account is the disappearance of the inverted region predicted by Equation 14.1 at large driving forces, because of the interference of the electronic states below the Fermi level that are thermodynamically unfavorable but kinetically advantageous. If necessary, nonadiabatic effects at the transition state could be included through modification of Z, for example, from the Landau–Zener theory [12,13] by introducing the transition probability p, the probability for charge transfer, being obtained from the electronic coupling energy C between the two involved states  2p  Z = χZ coll =   Z coll  (1 + p)    C 2 πRT  p = 1 − exp  −π    λ    RT 

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Dissociative Electron Transfers

O

O–

e–

+

O O

O

O 0

–2 4 –4

log (khet/cm s–1)

log (khom/M–1 s–1)

6

2 –6 0

–0.4

–0.8

–1.2 ΔG° (eV)

–1.6

FIgURE 14.3 Quadraticity of the activation–driving force relationship in the electrochemical (•) and homogeneous (○) reduction of pivaloyl peroxide. (Reprinted with permission from Antonello, S., Musumeci, M., Wayner, D.D.M., and Maran, F., J. Am. Chem. Soc., 123, 9577–9549. Copyright (2011) American Chemical Society.)

This set of equations were successfully applied to both heterogeneous and homogeneous CDETs (in the latter case, the electrode potential in the driving force expression should be replaced by the standard potential of the molecular electron donor). Many examples of the application of this model to the reduction kinetics of various families of compounds, involving C–halogen bonds (alkyl and benzyl halides) [6,14–16], O–O bonds (alkyl peroxides) [17,18], but also N–halogen bonds (N-halogeno sultams) [19], S–C bonds (sulfonium cations) [20], or S–Cl bonds in arene sulfenyl chlorides [21,22]. A remarkable example concerns the reduction of alkyl halides, for which ab initio calculations gave good evidence for applicability of the Morse curve model [23], while kinetics analysis of the homogeneous reaction between electrochemically generated aromatic anion radicals and tertiary alkyl halides leads to a very close agreement with theoretical calculated rate constants [15]. In the case of less hindered alkyl halides (secondary and primary), predicted rate constants are below experimental ones due to progressive competition between simple electron transfer and a SN2 pathway [24]. Another remarkable feature of the model derives from its quadratic character, as in the Marcus–Hush model. Large intrinsic barriers in CDET reactions lead to the fact that the curvature of the log k(E) versus ΔG0 is not easily observed. Combining homogeneous redox catalysis and direct electrochemical reduction of a family of peroxides with weak homolytic bond dissociation energies (DO–O ≈ 1.1–1.3 eV) allowed to investigate a large range of driving forces (about one and a half volt) and then unambiguously probe the curvature of the log rate constant versus driving force plot, as illustrated in Figure 14.3 for pivaloyl peroxide [25]. One additional interesting output of the CDET model is the determination of unknown homolytic bond dissociation energies from the experimental values of the reduction peak potentials obtained by cyclic voltammetry, providing that the concerted character of the reaction has been assessed and the standard potential of the leaving group is known. C–halogen, N–halogen, C–S, and O–O bond dissociation energies have thus been determined in various families of compounds.

B.

STICKY CONCERTED DISSOCIATIVE ELECTRON TRANSFER

During a dissociative electron transfer, an energy minimum, represented in Scheme 14.1 and in the potential energy diagram of Figure 14.1, corresponding to an ion–radical adduct (R•, X−) after

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Organic Electrochemistry

the cleavage has taken place, may be obtained. This adduct results from a charge–dipole attractive interaction between the cage fragments before they diffuse out, and it may be viewed alternatively as a σ* anion radical or as forming a weak three-electron bond. Even if these interactions strongly decrease from the gas phase to a polar liquid, they may partly survive when a partial positive charge is induced on the remaining radical part, thanks to the presence of a strong electron-withdrawing substituent [26,27]. The Morse curve model has been accordingly modified in order to take into account these effects. Potential energy curves describing both the reactant and the product systems are modeled by Morse curves and lead to a quadratic activation (activation free energy: ΔG≠)– driving force (standard free energy: ΔG 0) relationship as given in the following [26,27]:

∆G ≠ =

(

DRX − DR•, X− 4

)

2

 + λ0  1 +  

 ∆G − DR•, X−   2 DRX − DR•, X− + λ 0  

2

0

(

)

(14.2)

where D R•, X− is the interaction energy within the ion–radical pair. This last term, through its square root, has a marked influence of the electron transfer kinetics. As an example, if the sticky interaction amounts to ca. 1% of D RX, then a decrease of about 15% of the intrinsic barrier will ensue, thus accelerating the reaction to an experimentally detectable amount. To assess these sticky effects and quantify their magnitude, experiments in polar solvents have been done for families of compounds, on one hand by keeping the remaining radical constant while changing the leaving anion and on the other hand by keeping the leaving anion constant while changing the nature of the remaining radical. As concerns the former strategy, the electrochemical reduction in DMF of three haloacetonitriles (NCCH2X; X = Cl, Br, I) provides striking examples where the in-cage interaction rapidly decreases from Cl to Br and I (from about 40 meV to an almost vanishing interaction), thus showing a decreasing correlation with the halide radius [28]. Coming now to the latter strategy, it has been shown that the intensity of the interaction decreases as the polar character of the remaining radical decreases. Comparison within a family of polychloromethanes and polychloroethanes indeed indicates the following order of cluster interaction energies upon departure of a chloride anion, Cl3C–CCl3 (DR• ,Cl− ≈ 190 meV) > CCl4 > CHCl3, Cl2HC–CCl3 > CH2Cl2, Cl2HC–CHCl2, ClH2C–CHCl2 (DR• ,Cl− = 75 meV), in line with the expected inductive effects [29]. Increasing the solvation ability toward the leaving anion is still another way to probe these interactions. The 4-cyanobenzyl chloride reduction at an electrode in various solvents of increasing polarity leads to a decrease in the interaction energy within the (• CH2PhCN, Cl−) pair (from 135 meV in 1,2-dichloroethane to 40 meV in formamide), thus confirming the validity of the ad hoc sticky CDET model [27]. The model may also account for subtle intramolecular effects, for example, intramolecular hydrogen bonds, that may influence the dynamics of the cleavage reaction. As an example, it has been shown that with the 2-chloroacetamide (Scheme 14.2), the chloride anion interacts with the hydrogen atoms borne by the nitrogen upon reductive cleavage (DR•, Cl− ≈ 150 meV), and more strongly than it does with N,N-dimethylacetamide radical (obtained from 2-chloro-N,N-dimethylacetamide reductive dechlorination) in which the hydrogen atoms in the methyl groups are much less polarized (DR•, Cl− ≈ 100 meV) [30]. R R

H C Cl + e– H

N O

R N

H , Cl–

C

R

H O

R = H, Me

SCHEME 14.2

Sticky dissociative electron transfer to chloroacetamides.

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Dissociative Electron Transfers

As a final remark, one should note that the stepwise mechanism and the concerted should not be viewed as the extremes of a mechanism spectrum, according to the strength of the interaction between the two fragments. The stepwise and concerted pathways may enter in competition one with the other, and thus, the classical distinction between π and σ ion radicals seems more appropriate in this connection. The stepwise pathway involves a π ion radical that cleaves in an exergonic manner, thus giving rise to a σ ion radical, composed of weakly interacting fragments, finally yielding the separated fragments. The concerted pathway involves a σ ion radical that ultimately produces the separated fragments.

III.

COMPETITION bETWEEN CONCERTED AND SEQUENTIAL DISSOCIATIVE ELECTRON TRANSFERS

A.

STRUCTURAL, ELECTRONIC FACTORS AND SOLVENT EFFECTS CONTROLLING THE COMPETITION

The quadratic activation–driving force relationship that is obtained for both outer sphere electron transfer and CDET leads to the following expression for the transfer coefficient α: α=

∂∆G ≠ 1  ∆G 0 = 1 + ∂∆G 0 2  λ

 = 

∆G ≠ . λ

In contrast with outer sphere electron transfers, λ is significantly larger for CDET reactions, since it does include the bond dissociation energy of the cleaved bond. For a given scan rate, the activation free energy ΔG ≠ is almost the same for the two types of reactions; it thus follows that the transfer coefficient is much smaller in the concerted dissociative charge transfer. Determination of the transfer coefficient over a range of driving forces could be easily achieved through convolution techniques, and it could be also determined from peak width or variation of the peak width with the scan rate in cyclic voltammetry through linearization of the activation–driving force relationship around the peak value, thus providing a useful tool that helps discriminating between mechanisms [31]. α peak = − α=

RT F

  1.857    Epeak − Epeak/2 

29.5 (∂Epeak /∂ log v)

As illustrated by Figure 14.1, that depicts the reduction of a neutral substrate RX through a concerted or a stepwise pathway involving a radical anion RX as a transient (transposition to the cases where the reactant is charge positively or negatively is straightforward), the competition between the two routes involves thermodynamic and kinetic factors (internal factors), but also depends on the solvent and the energy of the incoming electron (external factors) as well as on electronic factors (coupling between the diabatic states involved). The passage from the stepwise to the concerted situation is expected to arise when the ion radical cleavage becomes faster and faster. Under these conditions, the rate-determining step of the stepwise process tends to become the initial electron transfer. Then, thermodynamics will favor one or the other mechanism according to the following equation:

(

0 0 0 = F − ERX/R + ERX/RX ∆Gcleav • •− +X −

(

)

0 = F D − EX0 • /X− + ERX/RX • − − T∆S RX →R• +X•

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)

(14.3)

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Organic Electrochemistry

0  is also the standard free energy of cleavage of the ion radical. Thus, one passes from the ∆Gcleav stepwise to the concerted mechanism as the driving force for cleaving the ion radical becomes larger and larger. It may thus be predicted that a weak bond dissociation energy (BDE) of the 0 R–X bond, a negative value of the standard potential of the transient (ERX/RX • −), and a positive 0 value of the standard potential of the leaving group (EX• /X−) will favor the concerted mechanism and vice versa. All three factors may vary from one RX molecule to another. However, families of compounds provide examples where the passage from one mechanism to the other is mainly driven by one of them, the changes in the others being minimized. Many studies have been conducted along these lines, and the influence of all of the factors has been identified and illustrated by several examples [16,19,32,33]. Some of them are reported in Table 14.1. In this context, the solvent plays an important role in the (de)stabilization of the intermediates as well as the leaving group involved in the dissociative charge transfer. For example, it has been shown that solvation and ion-pairing effects affect the cleavage reactivity of anion radicals containing a frangible bond, depending upon the localization of the negative charge [34]. For the anion radicals from 3-nitro-benzylchloride and bromide, 4-chlorobenzophenone, and N-fluoro-7-nitrosaccharin-sultam for which the negative charge is mostly concentrated on a small portion of the molecule (on the oxygen atom of carbonyl or nitro groups), the addition of water or ion-pairing agents (e.g., Li+ and Mg2+) slows down the cleavage of the anion radical, in parallel with a positive shift of the standard potential for its formation, that is, a stabilization of π ion radicals. In contrast, when the charge on the intermediate is spread out over the entire molecular framework, for example, in 2- and 9-chloroanthracenes, the addition of lithium or magnesium ions has no effect on the cleavage rate, while addition of water results in an accelerating effect at large water concentration caused by specific solvation of the leaving anion. Beyond these stabilization–destabilization effects, the solvent itself may be responsible for the passage from a stepwise to a concerted mechanism of the reductive cleavage reaction, that is, the very existence of π anion radicals may hinge on interactions between the electron first residence group and the solvent [35]. In the case of the para-cyanobenzyl chloride, gas phase calculations show that only a σ-radical anion at large C–Cl distances is obtained upon reduction [35], while reduction in water by pulse radiolysis goes through a radical anion intermediate, which decomposes quickly (submicrosecond time range) [36]. Calculations on a model compound including explicit solvent molecules (ONCH 2Cl + e− + 2H 2O) have allowed to capture the role of solvent molecules that stand close to the charge centers of the molecule, even if the representation of the solvent by only two water molecules is too simplistic at the quantitative level. During bond cleavage, solvent reorganization is well pictured in qualitative terms by the decrease in the interaction between one water molecule and the oxygen of the NO group and the parallel increase in the interaction between the second water molecule and the leaving chloride ion, as illustrated in Figure 14.4 [35]. The qualitative idea is that in a real solvent, the interactions of the many surrounding solvent molecules with the negative charge on the oxygen atom weaken at the benefit of the interactions between the solvent molecules surrounding the chlorine atom with the charge borne by this atom, the barrier for the anion radical cleavage being due to this solvent reorganization. In complement to the solvent effects on the thermodynamic and kinetic stability of π anion radicals and to the role of structural parameters, the electronic coupling between the diabatic states associated with the fragmented products on one hand and the intermediate on the other hand appears as an additional effect that may control the very existence of a transient species along the reaction pathway. Reductive cleavage of the three cyanobenzyl chloride isomers in N,N-dimethylformamide illustrates this idea [37]. While charge transfer is concerted with C–Cl bond breaking and acts as the rate-determining step in the case of both the ortho and para isomers, an anion radical is transiently formed before fragmentation for the meta isomer. In this family, the factors invoked so far to explain the followed mechanism are similar (bond dissociation

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TAbLE 14.1 Molecular Factors governing the Competition between Stepwise and Concerted Mechanisms Sequential

Concerted

Aryl halides, except some iodides, see text

All aliphatic halides

Examples of the prevailing role of E 0RX/RX CH2Cl (Br)

O2N

–

(NC) H

(NC) O2N

CH2Cl (Br)

N–F SO2

N–F SO2

O2N

C CH2Br O

References

(CH3O) H

C CH2Br O

[16] [19] [32]

Examples of the prevailing role of the bond dissociation energy (DRX) Cl (Br)

Z

Z

CH2Cl (Br)

[16]

(except Z = NO2)

NC

CH2F

NC

C CH2F O N–F SO2

O2N O2N

C CH2Cl (Br) O

O2N

CH2Cl (Br) O2N

Examples of the prevailing role of E 0X C CH2X + e– O X = Br, Cl

CH2Cl (Br)

N–Cl (Br, I) SO2

N–Cl (Br) SO2

[16, 32] [32, 33]

[19]

[16, 19]

/X –

C CH2 + X– O

[32]

X = OPh, OCH3, OC2H5, SPh, SC2H5, N(C2H5)2

energy, entropy of dissociation, standard potential of the leaving group, and in-cage interaction). As already mentioned, the electron transfer activation energy at the peak potential remains the same at a given scan rate; therefore, the potential peak shift observed between the meta isomer and the para and ortho isomers (about 200 mV more positive in the two latter cases) is a consequence of an energetic shift of the electron transfer transition state. Since it is very likely that the shape of the reactant potential energy curve is identical for all three isomers, the location change of the transition state is due to an energetic shift of the product electronic state curve. The reactant energy curve will respond by translating upward so that the activation energy

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Organic Electrochemistry 1.2 Potential energy (eV)

1 0.8 0.6 0.4 0.2 0 –0.2 –0.4

Reaction coordinate

–0.6 –0.8

–0.2 0 0.2 0.4 0.6 0.8 1 1.2 (a) 3.30

H O

N

2.40 Cl C

1.856 (b)

2.520

O

2.214

2.150

O

FIgURE 14.4 (See color insert.) Potential energy profiles (from QCISD(T)/6-31G* calculations) for cleavage of ONCH 2Cl anion radical in the presence of two water molecules. (a) Potential energy versus reaction coordinate. (b) Structures at each potential energy minimum (distances between atoms in Å). (Reprinted with permission from Costentin, C., Robert, M., and Savéant, J.-M., J. Am. Chem. Soc., 126, 2004, 16834–16840. Copyright (2011) American Chemical Society.)

remains identical. This vertical translation is endowed with a peak shift, that is, a change of the energy of the electron being transferred (see Figure 14.5a). The fundamental reason that underlie such shift is related to the fact that the π* orbital receiving the extra electron in the π anion radical − •RX does overlap with the σ* C–Cl orbital. In terms of electronic states, this means that the π* and σ* diabatic states mix. This mixing is expected to be heavily influenced by the nature and the position of substituents on the aromatic ring: it is stronger in the case of the ortho- and para-cyanobenzyl chlorides than it is in the case of the meta-cyanobenzyl chloride (the coupling constant in the phase space region where the two states' cross has been estimated to be larger than 0.15 eV in the former cases, in line with the fact that in the anion radical, the spin density is important at the para and ortho positions), so stronger that the π anion radicals no longer exist with the ortho and para compounds, that is, live less than a vibration, as illustrated in Figure 14.5b. In other words, the mechanistic pathway results from a balance between various effects such as intrinsic barrier for cleavage, solvent polarity (the more polar the solvent, the more stable the π anion radical), and also electronic coupling between the π* and σ* diabatic states, this coupling being modulated by the molecular structure. With the meta-cyanobenzyl chloride, the balance is in favor of the π anion radical existence in polar solvents such as acetonitrile, and the mechanism is driven toward a stepwise pathway. In the case of the para-(ortho-)cyanobenzyl chloride derivative, the balance may be reversed since a concerted mechanism occurs in 1,2-dichloromethane, acetonitrile, N,N-dimethylformamide, ethanol, and formamide, whereas a stepwise mechanism, with a very fast cleavage step, is observed in water, thanks to the solvent stabilizing effects on the intermediate anion radical [36].

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Dissociative Electron Transfers CH2Cl

CH2Cl

CN

CH2Cl

NC CN

Potential energy

RX

–E 0RX/RX – = –Ep,meta –Ep

Potential energy

– –

RX +

e–

R + X–

R + X– –E 0RX/R + X– (a)

RX

Reaction coordinate

Reaction coordinate (b)

FIgURE 14.5 (a) Potential energy profiles for the stepwise reduction of meta-cyanobenzyl chloride (dotted lines) and for the concerted reduction of ortho- and para-cyanobenzyl chlorides (full lines). (b) Diabatic (dotted lines) and adiabatic (full lines) electronic states of the products in the case of strong coupling (ortho and para derivatives).

B. ROLE OF THE DRIVING FORCE The electrochemical reduction of sulfonium cations in acetonitrile according to Figure 14.6 offers a 0 striking example of the combined roles of ERX/RX • − and of the homolytic bond dissociation energy, the former parameter being a measure of the energy of the π* orbital of the radical in which the incoming electron may be accommodated [20]. As seen in Figure 14.6, a matrix-type representation of the occurrence of the concerted and stepwise mechanism as a function of these two parameters may be sketched. The borderline situations that appear on the diagonal allow one to uncover the role of the driving force offered to the reaction. As mentioned earlier, an increase in the driving force offered to the reaction makes the mechanism pass from a concerted to a stepwise mechanism. The change is associated with a change of the activation driving force law, which remains quadratic but with different standard potentials and intrinsic barriers. The symmetry factors in the concerted mechanism are smaller, in absolute value, and vary less rapidly with the driving force than for the stepwise mechanism. In practice, cyclic voltammetry is an ideal tool to make the driving force vary, since the potential may be driven toward more negative values by increasing the scan rate or decreasing the temperature. A first experimental example of passage from a concerted to a stepwise mechanism was found with the reductive cleavage of the borderline sulfonium cations of Figure 14.6. An analysis of the observed variation of the transfer coefficient α with the scan rate shows a nonlinear variation, a characteristic of the transition zone between the two mechanisms [1,2,20]. Involvement of adsorption of the ylide formed during the reaction was shown to be insignificant, and thus, this unusual, nonlinear dependence of α with the driving force has been assigned to the transition between the two concerted (at small driving forces) and stepwise (at large driving forces) mechanisms. Another example of passage from one mechanism to the other was identified in the reduction of the tert-butyl 4-cyanoperbenzoate (O – O bond cleavage), for which cyclic voltammetry and convolution analysis have shown an “α-wave” with an almost constant symmetry factor at low driving forces, which progressively increases and then decreases when the driving force becomes larger and larger, the stepwise mechanism then becoming dominant [38]. Combination of

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Organic Electrochemistry E 0RX/RX

S+ S+– R + e– Concerted

S+

S+

CH3 Stepwise

S–R

–

CH3 Stepwise

CH3 Stepwise

Stepwise S+ S+R

S+ C H2

C H2 Borderline

D

Borderline

S+

S+ CN

C H2 Concerted

S+ C H2

CN

C H2 Concerted

CN

Stepwise

S+ C C H2 O Concerted

S+ C C H2 O Concerted

FIgURE 14.6 Reductive cleavage of sulfonium cations.

this cyclic voltammetry–convolution analysis also leads to the identification of a mechanism transition upon increasing the driving force in the heterogeneous dissociative reductive cleavage of S–C bonds in the 4-methylphenyl and 4-methoxyphenyl thiocyanate [39]. The electrochemical reduction of aryl iodides usually follows a stepwise mechanism, with a few exceptions. The mechanism shifts from concerted to stepwise upon increasing the scan rate in experiments carried out at room temperature with both iodobenzene and 4-methyliodobenzene (Figure 14.7). The transition between the two 0.45

0.45 α

α

0.4

0.4

0.35

0.35

log (v/V s–1)

log (v/V s–1)

0.3

0.3 –2

–1

0

1

–2

–1

0

1

2

FIgURE 14.7 Electrochemical reduction of aryl halides in DMF. Variation of the transfer coefficient α with the scan rate v. ●: iodobenzene, ○: bromobenzene, ▽: 1-iodonaphthalene, ◊: 4-methyliodobenzene, at 298 K, □: iodobenzene at 329 K.

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mechanisms is also obtained at higher temperatures, but the balance is, as expected, more in favor of the concerted mechanism. As compared to phenyl bromides, the C–halogen bond dissociation is weaker in iodide derivatives and their π* orbital less accessible than that of other aromatic substrates, for example, naphthalene iodides, for which the mechanism is always stepwise. Even if more difficult to detect, a mechanism crossover involving a homogeneous reaction has been characterized in the SRN1 reaction between the 2-nitropropane ion and the 4-nitrocumyl chloride [40]. In this case, the 2-nitropropanate ion serves both as nucleophile and electron inductor and the charge transfer from the 2-nitropropanate ion to the halo compound is faster by five orders of magnitude than predicted on the basis of a stepwise mechanism, thus leading to the conclusion that the concerted pathway is followed instead, due to a low driving force (+0.1 V). The diverse and accumulating examples thus gathered of a mechanistic shift from concerted to stepwise when increasing the driving force in the cleavage of C–Cl, C–I, C–S, and O–O bonds firmly confirm the key role of the energy of the incoming electron.

C.

ELECTROCHEMICAL VERSUS PHOTOINDUCED DISSOCIATIVE ELECTRON TRANSFER

Besides electrochemical induction (heterogeneous or homogeneous), reductive and oxidative cleavage reactions may also be triggered photochemically by means of an excited state of a sensitizer, as pictured in Figure 14.8. An interesting way of fighting energy-wasting back-electron transfer in photoinduced electron transfer reactions is to use a system where either the acceptor or the donor in the resulting ion pair, or both, undergoes a fast cleavage reaction [41]. The occurrence of a concerted dissociative reaction rather than a stepwise reaction (Figure 14.8a) thus appears as an extreme and ideal situation where the complete quenching fragmentation, quantum yield, should be unity [41]. From a diagnosis standpoint, the observation of a quantum yield smaller than unity would thus rule out the occurrence of a concerted dissociative mechanism. This intuitive notion is in fact incorrect as illustrated in Figure 14.8, showing a section of the potential energy surfaces along the reaction coordinate. After the photoinduced dissociative electron transfer has taken place, the system approaches the intersection between the fragmented product energy surface and the energy surface associated with the uphill dark dissociative electron transfer. There is thus the possibility of a partition of the reaction pathway between these two surfaces, leading in part to the fragmented products and in part to back-electron transfer. Figure 14.8 shows the case of an avoided crossing between the two surfaces, but partition may also occur in the case of a conical intersection. In any case, the quantum yield for a concerted dissociative photoinduced electron transfer may well be smaller than one [42]. Potential energy

D + hv D* + RX

D* D + RX Stepwise

(R , X , D*)

(R , X , D) (RX, D*) hv

(R , X–, D +)

Concerted R + X– + D (a)

(RX, D) (b)

Reaction coordinate

FIgURE 14.8 Photoinduced dissociative electron transfer. Reaction scheme (a). Section of the zero-order (...) and first-order (_) potential energy surfaces along the reaction coordinate, when stretching of the cleaving bond is the main factor of nuclei reorganization (b).

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Organic Electrochemistry CN CN

CF3

S+

CCl4 CH2

S +

CH2Cl

S +

CH2 CN

(a)

C

C

C

Borderline

S

S

(b)

C

C

Borderline

S

S

S

(C: concerted reductive cleavage, S: sequential reductive cleavage)

SCHEME 14.3 Similarities and differences in the mechanisms of (a) electrochemical and (b) photoinduced (with the excited state of 2-ethyl-9,10-dimethoxyanthracene as an electron donor) reductive cleavage of various compounds.

For the same substrate, the driving force offered in the photoinduced case is always larger than in the electrochemical case, making it possible for the mechanism to be stepwise or borderline in the first case and borderline or concerted in the second [43]. As an example, the reductive cleavage of the C–S bond in 4-cyanobenzylmethylphenyl sulfonium cation follows a concerted pathway in electrochemical conditions, while its photoinduced reduction by the excited state of 2-ethyl-9,10-dimethoxyanthracene appears to be a competition between the concerted and the stepwise pathways, both being equally favored [43]. Scheme 14.3 highlights the differences and similarities between electrochemically and photochemically induced dissociative electron transfers in a series of substrates. The factors that control the competition between the two mechanisms electrochemical and photoinduced electron transfer are thus coherent, provided the differences in driving force are taken into account.

IV. CLEAVAgE OF PRIMARy RADICALS Primary radicals, for example, ion radicals, issued from one electron transfer to or from a parent molecule RX are usually fragile species that quickly decompose, more readily than their parents. When the unpaired electron is located on the R portion of the molecule, cleavage is heterolytic and amounts to an intramolecular electron transfer from the R to the X leaving group in the case of an anion radical and to an intramolecular electron transfer in the reverse direction in the case of a cation radical (Scheme 14.4). When on contrary the unpaired electron is located on the X portion of the molecule, the cleavage step involves a homolytic dissociation of the fragmented bond (Scheme 14.4). This distinction between two modes of cleavage is also valid in the case of uncharged radicals resulting either from the reduction of cationic substrates or to the oxidation of anion substrates, producing a neutral secondary radical, R− •, and a neutral leaving group, X. In both the homolytic and heterolytic case, it has been shown that the cleavage reactions could be described as an intramolecular dissociative electron transfer concerted with bond breaking [44,45].

–

R

X

R

R +

SCHEME 14.4

+

–

Homolytic

Heterolytic (a)

X

R• X

Heterolytic

R

X

+

Homolytic R + X–

X– (b)

Heterolytic and homolytic bond cleavages in radical anions (a) and radical cations (b).

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Dissociative Electron Transfers

The cleavage of aromatic anion radicals, which provides an example of heterolytic cleavage, has attracted considerable attention, with emphasis on the cases where the leaving group is a halide ion, stimulated by the fact that this type of reaction is involved in the propagation loop of electron transfer–catalyzed aromatic SRN1 substitutions. An extension of the theory of dissociative electron transfer has been proposed to the intramolecular case, leading to the following quadratic activation– driving force relationship [44,46]: ≠ RX• − →R• + X −

∆G

0  ∆GRX •− →R• + X − = ∆G  1 +  4∆G0≠  ≠ 0

  

2

(14.4)

where ≠ ∆GRX • − →R• + X −  is the free energy of activation

0 ∆GRX • −→R• + X −  is the standard free energy for cleavage and is identical to Equation 14.3

(

0 0 0 0 ∆GRX • − →R• + X − = ∆Gcleav = F DRX →R• + X• − T∆SRX →R• + X• + ERX/RX• − − EX• /X −

)

where D RX→R• + X• is the homolytic bond dissociation energy of the starting RX ΔSRX→R• + X• is the corresponding entropy change The E 0 are the standard potentials of the various subscript redox couples The intrinsic barrier, ∆G0≠ , is equal to one-fourth of the sum of the homolytic bond dissociation energy of the intermediate and of the reorganization energy upon cleavage: ∆G0≠ =

DRX• − + λ 0

(14.5)

4

where DRX•− may be expressed as 0 0 DRX• − = DRX→R• + X• + ERX/RX • − − ER• /[R• ]• −

(the two missing entropic terms, T (SRX − SRX• −) and −T(SR•−S[R•]• −), may be considered as compen• • sating each other) and [R ] − represents a species obtained from the injection of one electron in the • π* orbital of the σ-radical R , thus leading to an excited state of the carbanion R−. λ0, the solvent reorganization energy, corresponds to the transfer of the negative charge from the anion radical to the leaving halide ion, and it could be expressed in the Marcus way as

λ0 =

e02  1 1  1 1 1  −  + − 4πε0  εop εs   2aRX• − 2aX− aRX• − + aX−  

   

where e0 is the electron charge ε0 is the vacuum permittivity εop and εS are the solvent optical and static dielectric constants, respectively The a are the radii of the equivalent spheres of the subscript species.

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Organic Electrochemistry 12

12

log k

10

8 6

4

4

2

2 0 –2 –4

0 –2 –4

0 – –ERCl/RCl

0.8

1.2

log k

10

8 6

1.6

2

2.4

2.8

1.2

0 –ERBr/RBr –

0.6

1

1.4

1.8

2.2

2.6

1.2 ΔG≠

1

RCl

–

–

R + Cl

0.8

0.8

0.6

0.6

0.4

0.4

0.2 ΔG0

0

RCl

ΔG≠

1

RBr –

R + Br–

0.2 –

R + Cl–

–0.2

ΔG0

0

RBr –

R + Br–

–0.2 1.5

1

0.5

0

–0.5 –1

1.5

1

0.5

0

–0.5 –1

FIgURE 14.9 Correlation between the fragmentation rate constant (in s−1) and the standard potential (in V vs. SCE) (top) and activation–driving force relationship (free enthalpies in eV) (bottom) for aryl chloride and bromide anion radicals. ≠ The activation free energy, ∆GRX • − →R• + X − , may also be obtained from the experimental determiRT ≠ nation of the cleavage rate constant k: ∆GRX ln(kBT /hk ) . Combining electrochemical •− →R• + X − = F and photochemical techniques (e.g., pulse radiolysis) has allowed to determine cleavage rate constants spanning more than 12 orders of magnitude, for both chloro and bromo aromatic compounds 0 [46]. The activation-standard potential plots (ERX/RX •−) for both aryl chlorides (23 compounds) and aryl bromides (22 compounds) show remarkable correlation (Figure 14.9). Converting then the 0 ERX/RX •− scale into a driving force scale through the use of Equation 14.3 leads to activation–driving force plots shown in Figure 14.9. These plots necessitate the estimation of both EX0 • /X− , the bond dissociation energies D RX→R• + X• of the neutral compounds (almost constant for a given halogen) and the entropic term (the quasiconsistency of which were checked by DFT calculations). In the two cases, the correlation straight line is close to 0.5, thus suggesting, taking into account that the driving force range is spread out over positive and negative values, a linearization of the quadratic activation–driving force relationship (14.4) [46]:

1 ≠ 0 ∆GRX = ∆G0≠ + ∆GRX •− •− →R• + X − →R• + X − 2

(14.6)

The intrinsic barrier, ∆G0≠, may thus be estimated to 0.41 eV for the chlorides and to 0.39 eV for the bromides. Application of Equation 14.5 leads to a predicted estimation of ∆G0≠ , 1.23 and 1.03 eV for the chloro and bromo derivatives, respectively. The large difference obtained between the predicted and experimental values may originate from two factors that tend to decrease the barrier to overcome for cleavage. The first one is out-of-plane bending that permits skirting round

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Dissociative Electron Transfers

the conical intersection encountered upon straight stretching of the C–X bond as in the model used to derive Equation 14.4. The second one is due to the possible formation of a weak σ anion radical after cleavage has taken place (D R•,X−), as described in Section II.B for purely concerted processes. The previous Morse curve model was adapted to take into account these two additional factors [46]. The first of these is assumed to follow a harmonic law, that is, the energy increase varies as ( fb /2)θ2b , where θb is the bending angle and f b the force constant. The gain in resonance energy is assumed to be proportional to the bending angle: H = h 0 θb. It follows that Equation 14.4 remains valid but with a different definition of the intrinsic barrier:

λ0 + ∆G0≠

=

(

2

DRX• − − DR• ,X−

) −h

2 0

(14.7)

2 fb

4

h02 / 2 fb may be regarded as constant within the linearized region around the zero of driving force. The activation–driving force relationship (14.6) is thus modified:

λ0 + ≠ ∆GRX • − →R• + X − =

(

2

DRX• − − DR• ,X− 4

0 + 0.5∆GRX • − →R• + X −

) −h

2 0

2 fb (14.8)

It thus appears that the intrinsic barrier is decreased through two terms, −h02 / 2 fb and DR• ,X− . Values of −h02 / 2 fb comprised between 0.3 and 0.5 eV, as computed for 4-cyanochlorobenzene are compatible with the aforementioned experimental data. The reduction of the 2-nitro and 2,4-dinitrophenyl sulfenyl chloride provides another example where the initially formed π radical anion upon one electron reduction cleaves heterolytically and yields a σ* radical anion before the sulfur-centered radical and the chloride anion fall apart [22]. The model describing the heterolytic cleavage of anion radicals also proved to be useful in deciphering regioselectivity of the cleavage, for example, for the reduction of nitro-substituted benzyl thiocyanates (O2N–C6H4CH 2 –S–CN), for which α-cleavage (CH 2 –S bond) is preferred over β-cleavage (S–CN bond) from both thermodynamic and kinetic points of view [47]. Homolytic, exothermic cleavage of radicals may exhibit a sizable activation barrier, for example, in the electrochemical reduction of the 4-nitro-tert-butylperbenzoate, whose O–O bond is broken after the formation of a transient π* anion radical [45]. Among various examples, this is also the case in the cleavage of anion radicals of nitrobenzyl phenyl ether (C–O bond cleavage), as well as during the cleavage of the S–C bond in the anion radical of the nitrophenyl diphenylmethyl thioether, and of the C–N bond in the anion radical of α-nitrocumyl. Such homolytic cleavage of radicals and ion radicals implies that an electron is being transferred from a π* orbital into the σ* orbital of the bond being broken. It has been shown, based on a two-state description model of the potential energy surfaces, and inclusion of the bond cleavage and formation as Morse curves in the normal-mode analysis, that the kinetics of cleavage could be described through a quadratic activation–driving force relationship [45]. The intrinsic barrier is made of two main contributions, the π* to σ* excitation energy in the transient ion radial and the triplet excitation energy of the leaving group. The predictions of the model in terms of fragmentation kinetics agree well with experiments in a family of 4-cyanophenyl alkyl ether anion radicals [45].

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Organic Electrochemistry

V. CONCLUDINg REMARKS Activation–driving force relationship based on Morse curve description of potential energies of the various electronic states has proved their validity for both qualitative and quantitative descriptions of dissociative electron transfers, including concerted reactions, homolytic and heterolytic cleavage of primary radicals and ion radicals, as well as photoinduced processes, for which the electron donor is the excited state of an aromatic molecule. Backed with experimental facts, the understanding of the coupling of bond breaking with electron transfer has now reached a good degree of comprehensiveness, in terms of kinetics, mechanisms, and models. The factors that govern the dichotomy and competition between stepwise and concerted pathways, encompassing structural and electronic factors as well as the influence of the solvent and the energy of the incoming electron, have been deciphered and their role analyzed. Recent improvements lead to modeling modification to take into account the possible attractive interactions between caged fragments (sticky dissociative electron transfer) and symmetry restrictions leading to conical intersections. A route is therefore opened to widespread applications of the available concepts and models, with possible emphasis on biological processes, especially redox enzyme reactions. As an example, it is worth mentioning that reductive dehalogenases isolated from anaerobic bacteria, which couple the reductive dehalogenation of tetrachloroethylene (PCE), trichloroethylene (TCE), and other chlorinated hydrocarbons to energy conservation (dehalo-respiration), open the way to biological strategies for remediation of systems polluted by halohydrocarbons. Sulfurospirillum multivorans catalyzes the reductive dehalogenation of PCE and TCE to (Z)-1,2-dichloroethene (DCE) with exceptionally high specific activities, using as cofactor a corrinoid (norpseudovitamin B12) [5]. The rate constants for electron transfer from the enzyme or from outer sphere aromatic reduced donors to PCE or TCE were derived from cyclic voltammetric redox catalysis experiments, and were slower with aromatic donors by 12 orders of magnitude [48]. An electron transfer mechanism is thus not consistent with the observed rates. The actual mechanism therefore involves more intimate interactions between the electron donor and the substrate in which the PCE (or TCE) molecule enters the cobalt coordination sphere. Beyond this example, there are clearly a large number of other enzymatic processes in which electron transfer and bond breaking are coupled, the understanding of which could benefit from the same kind of electrochemical approach. Electron transfer–/bond breaking–coupled reactions illustrate the benefits that could result from concertedness. It results in an increase of the driving force offered to the overall process; however, there is a price to pay, due to the increase in the intrinsic barrier paralleling the increase in the bond strength, even if the global result is beneficial. Coupling of electron transfer and proton transfer is another example of the virtues of concertedness. Coupling between electron transfer and proton transfer indeed occurs in a huge number of natural and artificial processes, for example, in the electron transfer activation of small molecules such as H2O, O2, and CO2, and the growing attention that these reactions focuses is likely to intensify by the necessity to face current energy challenges. Focusing on reactions where proton and electron transfers involve different sites, the possibility that proton and electron transfer steps are concerted giving rise to concerted proton–electron transfer (CPET) reactions as opposed to stepwise pathways in which proton transfer precedes (PET) or follows (EPT) electron transfer, possess the advantage of circumventing the stepwise pathway high-energy intermediates. This thermodynamic benefit is, however, associated with a kinetic counterpart related to an increased nonadiabaticity due to the tunneling of both electron and proton in the transition state rather than to an increased reorganization energy (see Chapter 13). The next step would be to see if concertedness could apply to reactions involving not only electron transfer and proton transfer but also breaking of a bond between heavy atoms. The impact of concertedness for such couplings would be particularly worth investigating in the context of small molecule activation, for example, in the electrochemical reduction of CO2, whose first step consists in the formation of CO, which implies a two-electron reduction, a C–O bond cleavage, and a protonation of the leaving oxygen atom. Only one example of an all-concerted

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529

reaction has been shown in the reductive cleavage of an O–O bond [49], helped by the presence of a proximal, H-bonded carboxylic acid group, the experimental and theoretical analysis of which is detailed in Chapter 13.

REFERENCES 1. Houman, A. Chem. Rev. 2008, 108, 2180–2237. 2. Costentin, C.; Robert, M.; Savéant, J.-M. Chem. Phys. 2006, 324, 40–56. 3. Savéant, J.-M. Electron transfer, bond breaking and bond formation. In: Advances in Physical Organic Chemistry, Tidwell, T. T. (Ed.), Academic Press: New York, 2000, vol. 35, pp. 117–192. 4. Maran, F.; Wayner, D. D. M.; Workentin M. S. In: Advances in Physical Organic Chemistry, Tidwell, T. T. (Ed.), Academic Press: New York, 2001, vol. 36, pp. 85–116. 5. Neumann, A.; Scholz-Muramatsu, H.; Diekert, G. Arch. Microbiol. 1994, 162, 295–301. 6. Savéant, J.-M. J. Am. Chem. Soc. 1987, 109, 6788–6795. 7. Marcus, R. A. J. Chem. Phys. 1956, 24, 966–978. 8. Hush, N. S. J. Chem. Phys. 1958, 28, 962–972. 9. Marcus, R. A. Electrochim. Acta 1968, 13, 995–1004. 10. Hush, N. S. Electrochim. Acta 1968, 13, 1005–1023. 11. Marcus, R. A. J. Chem. Phys. 1965, 43, 679–701. 12. Landau, L. Phys. Z. Sowjetunion 1932, 2, 46–51. 13. Zener, C. Proc. R. Soc. Lond. Ser. A 1932, 137, 696–702. 14. Clark, K. B.; Wayner, D. D. M. J. Am. Chem. Soc. 1991, 113, 9363–9365. 15. Savéant, J.-M. J. Am. Chem. Soc. 1992, 114, 10595–10602. 16. Andrieux, C. P.; Le Gorande, A.; Savéant, J.-M. J. Am. Chem. Soc. 1992, 114, 6892–6904. 17. Workentin, M. S.; Maran, F.; Wayner, D. D. M. J. Am. Chem. Soc. 1995, 117, 2120–2121. 18. Antonello, S.; Musumeci, M.; Wayner, D. D. M.; Maran, F. J. Am. Chem. Soc. 1997, 119, 9541–9549. 19. Andrieux, C. P.; Differding, E.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 1993, 115, 6592–6599. 20. Andrieux, C. P.; Robert, M.; Saeva, F. D.; Savéant, J.-M. J. Am. Chem. Soc. 1994, 116, 7864–7871. 21. Ji, C.; Goddard, J. D.; Houmam, A. J. Am. Chem. Soc. 2004, 126, 8076–8077. 22. Ji, C.; Ahmida, M.; Chahma, M.; Houmam, A. J. Am. Chem. Soc. 2006, 128, 15423–15431. 23. Bertran, J.; Galardo, I.; Moreno, M.; Savéant, J.-M. J. Am. Chem. Soc. 1992, 114, 9576–9583. 24. Costentin, C.; Savéant. J.-M. J. Am. Chem. Soc. 2000, 122, 2329–2338 and references therein. 25. Antonello, S.; Formaggio, F.; Moretto, A.; Toniolo, C.; Maran, F. J. Am. Chem. Soc. 2001, 123, 9577–9584. 26. Pause, L.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 2000, 122, 9829–9835. 27. Pause, L.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 2001, 123, 11908–11916. 28. Cardinale, A.; Isse, A. A.; Gennaro, A.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 2002, 124, 13533–13539. 29. Costentin, C.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 2003, 125, 10729–10739. 30. Costentin, C.; Louault, C.; Robert, M.; Teillout, A.-L. J. Phys. Chem. A 2005, 109, 2984–2990. 31. Savéant, J.-M. Elements of Molecular and Biomolecular Electrochemistry, Wiley-Interscience: New York, 2006. 32. Andrieux, C. P.; Savéant, J.-M.; Tallec, A.; Tardivel, R.; Tardy, C. J. Am. Chem. Soc. 1997, 119, 2420–2429. 33. Andrieux, C. P.; Combellas, C.; Kanoufi, F.; Savéant, J.-M.; Thiébault, A. J. Am. Chem. Soc. 1997, 119, 9527–9540. 34. Andrieux, C. P.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 1995, 117, 9340–9346. 35. Costentin, C.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 2004, 126, 16834–16840. 36. Neta, P.; Behar, D. J. Am. Chem. Soc. 1981, 103, 103–106. 37. Costentin, C.; Donati, L.; Robert, M. Chem. Eur. J. 2009, 15, 785–792. 38. Antonello, S.; Maran, F. J. Am. Chem. Soc. 1999, 121, 9668–9676. 39. Houmam, A.; Hamed, E. M.; Still, I. W. J. J. Am. Chem. Soc. 2003, 125, 7258–7265. 40. Costentin, C.; Hapiot, P.; Médebielle, M.; Savéant, J.-M. J. Am. Chem. Soc. 1999, 121, 4451–4460. 41. (a) Saeva, F. D. Top. Curr. Chem. 1990, 156, 59–92; (b) Gaillard, E. R.; Whitten, D. G. Acc. Chem. Res. 1996, 29, 292–297. 42. Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 2000, 122, 514–517. 43. Pause, L.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 2001, 123, 4886–4895.

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Savéant, J.-M. J. Phys. Chem. 1994, 98, 3716–3734. Costentin, C.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 2003, 125, 105–112. Costentin, C.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 2004, 126, 16051–16057. Hamed, E. M.; Doai, H.; McLaughlin, C. K.; Houmam, A. J. Am. Chem. Soc. 2006, 128, 6595–6604. Dieckert, G.; Gugova, D.; Limoges, B.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 2005, 127, 13583–13588. 49. Costentin, C.; Hajj, V.; Robert, M.; Savéant, J.-M.; Tard, C. Proc. Natl. Acad. Sci. USA 2011, 108, 8559–8564.

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15

Electron Transfer–Catalyzed Reactions Kazuhiro Chiba and Yohei Okada

CONTENTS I. II.

Introduction ........................................................................................................................... 531 Oxidative Mediators .............................................................................................................. 534 A. General Mechanism ....................................................................................................... 534 B. Triarylamines................................................................................................................. 534 C. TEMPO ......................................................................................................................... 538 D. Halide Ions ..................................................................................................................... 538 E. Iodoarenes ...................................................................................................................... 539 F. o-Aminophenols ............................................................................................................ 543 III. Reductive Mediators..............................................................................................................544 A. General Mechanism .......................................................................................................544 B. Nickel(II) Salens ............................................................................................................544 C. o-Carboranes ................................................................................................................. 545 IV. Oxidative Chain Reactions ....................................................................................................546 A. General Mechanism .......................................................................................................546 B. Radical Cation Chain Reaction ..................................................................................... 547 C. Cation Chain Reaction ................................................................................................... 547 V. Reductive Chain Reactions ................................................................................................... 549 A. General Mechanism ....................................................................................................... 549 B. Electrochemical SRN1 Reaction ..................................................................................... 550 VI. Measuring Redox Potentials.................................................................................................. 551 A. General Aspects ............................................................................................................. 551 B. Direct Measuring ........................................................................................................... 551 C. Indirect Measuring ........................................................................................................ 551 VII. Conclusions ........................................................................................................................... 552 References ...................................................................................................................................... 553

I. INTRODUCTION Electron transfer is one of the most fundamental and ubiquitous processes of all chemical and biological systems, and thus has been studied in detail both theoretically and practically. In organic synthesis, electron transfer–induced reactions have been extensively utilized to achieve various chemical transformations, constructing a wide variety of organic compounds, including natural products, pharmaceutical products, and functional materials. In order to trigger electron transfer– induced reactions, photochemical processes and one-electron redox reagents are widely employed. In this context, electrochemical processes have also been utilized to regulate either one or two electron transfers at the surface of the electrodes that afford not only various functional group transformations but also a wide variety of carbon–carbon bond formation reactions in a heterogeneous manner [1–3]. Several reactive intermediates such as ions, radicals, and radical ions are 531

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Organic Electrochemistry

generated through the electron transfer to avoid consumption of additional redox reagents and the redox potentials can be easily controlled under mild electrolytic conditions. Therefore, they are promising methodologies from an environmental viewpoint. Moreover, various elegant strategies have been demonstrated to modify the surface of the electrodes, which might dramatically change the reaction conditions. It is well established that gold or carbon electrodes can be covalently grafted through treatment with thiols [4–13] or aryldiazonium salts [14–23] to functionalize their surfaces. For example, a self-assembled monolayer (SAM) was formed enantioselectively on the surface of a gold electrode with a (111)-oriented surface using d- or l-homocysteine (Hcy) (Scheme 15.1) [24]. The redox behavior of an electrochemically active chiral 3,4-dihydroxyphenylalanine (DOPA) was then analyzed by cyclic voltammetry using the Hcy-modified gold electrode to impart enantioselectivity in the redox reaction of DOPA in acidic solution. Thus, the SAM of d-Hcy (1) was suggested to block the redox reaction of l-DOPA (2), while the SAM of l-Hcy (3) was suggested to block that of d-DOPA (4). These modified electrodes have found a wide variety of applications in chemistry, including analytical, biological, and physical. In contrast, it is well known that electron transfer between two solid phases is severely limited. In organic electrochemistry, solid phase–bound substrates are barely oxidized or reduced at the surface of the electrodes. Based on this fact, solid phase–supported bases have been introduced into electrolytic systems for in situ generation of a supporting electrolyte [25–28]. For example, methanol was NH+3

NH+3 O

HS 1

OH

3

NH+3

OH

NH+3 O

O

S

O

HS

OH

S HO

HO

OH

NH+3

O

O

HO

O

HO

OH

2

O O

O

OH

2

NH+3

NH+3 O

O

OH

NH+3

NH+3

OH

NH+3 O

O S

S OH

HO

OH

NH+3

HO

NH+3 O

O HO

HO OH

4 O

OH

4 O

NH+3

NH+3 O

O O

O OH

NH+3 O S

S

O OH

OH Gold electrode

Gold electrode

SCHEME 15.1

OH

NH+3

Enantioselective redox reaction of DOPA using homocysteine modified gold electrodes.

© 2016 by Taylor & Francis Group, LLC

533

Electron Transfer–Catalyzed Reactions Base

S

CF3

5 S OMe S

CF3

MeOH Base

CF3

5

10 mA/cm2 3F MeOH (+)Pt-Pt(–)

Solid phase

OMe

Base

Porous polystyrene Silica gel

CF3

S

Yield (%)

92

N

76

N

MeO– + H+ Silica gel

89

N

Anode

NH

SCHEME 15.2 Anodic methoxylation of phenyl(2,2,2-trifluoroethyl)sulfide (5) using solid-supported bases. 

dissociated into methoxide anions and protons by porous polystyrene or silica gel–supported bases to achieve electrical conductivity without additional supporting electrolytes, while solid phase–supported bases were not oxidized at the surface of the electrodes (Scheme 15.2) [29]. The anodic methoxylation of phenyl(2,2,2-trifluoroethyl)sulfide (5) took place effectively in the presence of solid phase–supported bases. In addition, these solid phase–supported bases could be recovered simply by filtration, enabling their reuse. Surface-bound and immobilized molecules are discussed in detail in Chapter 42. As described, in general, direct heterogeneous electron transfer at the surface of electrodes induces electrochemical transformations. Meanwhile, several reactive intermediates generated through the electron transfer can react in a homogeneous bulk electrolyte solution. Therefore, the possibility of organic electrochemistry can be expanded through the use of indirect homogeneous transformations. For example, cathodic reduction of the starting material results in electrochemical reactions forming anionic species that function as subsequent bases or nucleophiles, known as electrogenerated bases and nucleophiles, which are discussed in detail in Chapter 43. In this chapter, indirect electrochemical transformations using redox mediators are discussed (Scheme 15.3). The redox mediators are recognized as electron carriers, which are activated through heterogeneous electron transfer at the surface Homogeneous electron transfer

Heterogeneous electron transfer Mediator (inactive)

Starting substrate

Mediator (active)

Reactive intermediate

Chemical transformations

Electrode

SCHEME 15.3

Schematic illustration of an indirect electrochemical transformation using redox mediators.

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534

Organic Electrochemistry Homogeneous electron transfer

Heterogeneous electron transfer

Starting substrate

Reactive intermediate

Chemical transformations

Reactive intermediate

Electrode

SCHEME 15.4

Schematic illustration of an electrochemical chain reaction.

of electrodes and then function as homogeneous redox reagents. One of the most salient features of indirect electrochemical transformations using redox mediators is that they take place over the existence of thermodynamic hurdles that would seemingly preclude electron transfer between the activated mediators and substrates in bulk electrolyte solution [30,31]. After the homogeneous electron transfer, the redox mediators are generally reactivated at the surface of the electrodes through heterogeneous electron transfer. Therefore, only a small amount of the redox mediators is required to complete the chemical transformations. To date, various organic and inorganic mediators have been established to realize the ingenious design of the reaction systems, leading to promising manufacturing techniques. Additionally, autoactivating chemical transformations induced by direct heterogeneous electron transfer, known as electrochemical chain reactions, are discussed in this chapter (Scheme 15.4). In this case, the starting substrate is activated through electron transfer at the surface of the electrode, while the resulting intermediate is responsible for the following activation in the homogeneous bulk electrolyte solution to drive the chain cycle. Thus, only a catalytic amount of electricity is required to complete the reaction. The aim of this chapter is to give an outline of indirect electrochemical transformations with recent examples.

II.

OXIDATIVE MEDIATORS

A.

GENERAL MECHANISM

As indicated earlier, redox mediators can be defined as electron carriers that are activated through heterogeneous electron transfer at the surface of electrodes and then function as homogeneous redox reagents. There are two types of redox mediators: oxidative and reductive. Typically, the oxidative mediators are activated through anodic oxidation and the resulting intermediates function as subsequent oxidants (Scheme 15.5). The oxidative mediators also play critical roles in preventing overoxidation of the products, which suppresses the formation of a polymeric film on the anode, known as anodic passivation [32–34]. In this section, indirect electrochemical transformations using oxidative mediators, including triarylamines, 2,2,6,6-tetramethylpiperidine-1-oxyl (TEMPO), halide ions, iodoarenes, and o-aminophenols, are discussed with their use in recent synthetic examples.

B.

TRIARYLAMINES

It has been well established that the radical cations of triarylamines are highly stable; thus, they can be utilized as one-electron oxidants in chemical transformations [35,36]. For example, trans-anethole (6) was oxidized by the radical cation of tris(4-bromophenyl)amine (7), which is commercially

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535

Electron Transfer–Catalyzed Reactions Bulk oxidation

Anodic oxidation Mediator (reduced)

Substrate (reduced)

–e–

+e– –e– Mediator (oxidized)

Substrate (oxidized)

Chemical transformations

Anode

SCHEME 15.5

Schematic illustration of an indirect anodic transformation using oxidative mediators. MeO MeO

7

Br

Br

+SbCl– 6

N

6 Br 7

MeO Br

Br N

MeO 6

Br 7

+e– –e–

MeO

MeO

Br

Br +

MeO

6

N + Br

MeO

SCHEME 15.6 Oxidative dimerization of trans-anethol (6) using tris(4-bromophenyl)amine (7) as an oxidative mediator.

available as a tris(4-bromophenyl)aminium hexachloroantimonate, to give the corresponding radical cation, leading to dimerization (Scheme 15.6) [37]. The mechanism of the dimerization has been discussed in detail [38–40]. Meanwhile, such radical cations of the triarylamines can also be obtained through the anodic oxidation of the corresponding triarylamines [41]. After the bulk oxidation, the neutral triarylamines are reoxidized at the surface of the anode; thus, they function as oxidative mediators. In these cases, the oxidation potentials of the triarylamines can be modulated as a function of different substitution patterns of the aromatic rings (Table 15.1) [42]. With these triarylamines in hand, tris(4-methyl-2-nitrophenyl)amine (8) was then employed as an oxidative mediator to demonstrate the indirect electrochemical cleavage of electron-deficient substituted stilbenes [43]. These stilbenes were oxidized to give the corresponding radical cations, which were then trapped by water, leading to the electrochemical equivalent of ozonolysis to afford the corresponding aldehydes (Scheme 15.7) [44]. Because the oxidation potentials of these stilbenes are close to that of the solvent, direct electrochemical cleavages required a large excess of electricity to consume the starting substrates, which also caused undesired side reactions attributed mainly to the oxidation of solvent. In contrast, indirect electrochemical cleavage using 8 as an oxidative mediator required relatively small amounts of electricity to complete the reactions (Scheme 15.8). In these cases, the

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536

Organic Electrochemistry

TAbLE 15.1 Oxidation Potentials of Triarylamines R1

R1 R6

R7

R2

+

vs Ag/AgNO3

N R5

R3

R6

R7

R2

N

R5

R3

0.1 M LiClO4 CH3CN (+)C–Pt(–)

R4

R4

Oxidation potentials (vs Ag/AgNO3)

Substrates R1 = H, R2 = OMe, R3 = H, R4 = OMe, R5 = H, R6 = OMe, R7 = H

E OX = 0.25 V

R1 = H, R2 = Me, R3 = H, R4 =

E OX = 0.40 V

Me, R5 = H, R6 = Me, R7 = H

R1 = H, R2 = H, R3 = H, R4 = H, R5 = H, R6 = H, R7 = H

E OX = 0.54 V

7 (R1 = H, R2 = Br, R3 = H, R4 = Br, R5 = H, R6 = Br, R7 = H)

E OX = 0.78 V

R1 = H, R2 = Cl, R3 = H, R4 = Cl, R5 = H, R6 = Cl, R7 = H

E OX = 0.79 V

R1 = H, R2 = H, R3 = NO2, R4 = H, R5 = H, R6 = H, R7 = H

E OX = 0.83 V

R1 = NO

E OX = 0.84 V

, R2 = H, R3 = H, R4 = H, R5 = H, R6 = H, R7 = H

2 R1 = H, R2 = NO2, R3 = H, R4 = H, R5 = H, R6 = H, R7 = H

R1 = H, R2 = Me, R3 = NO

, R4 = Me, R5 = NO

2

E OX = 0.88 V

6 7 2, R = Me, R = H

E OX = 0.96 V

R1 = H, R2 = Cl, R3 = NO2, R4 = Cl, R5 = H, R6 = Cl, R7 = H

E OX = 1.00 V

R1 = H, R2 = NO2, R3 = H, R4 = NO2, R5 = H, R6 = H, R7 = H

E OX = 1.08V

R1 = H, R2 = NO2, R3 = H, R4 = Br, R5 = NO2, R6 = Br, R7 = H

E OX = 1.17 V

R1 = H, R2 = Cl, R3 = NO2, R4 = Cl, R5 = NO2, R6 = Cl, R7 = H

E OX = 1.25 V

8 (R1 = H, R2 = Me, R3 = NO

E OX = 1.28 V

R1 = H, R2 = NO2, R3 = H, R4 = Br, R5 = Br, R6 = NO2, R7 = H

E OX = 1.31 V

4 5 6 7 2, R = Me, R = NO2, R = Me, R = NO2)

1 = H, R2 = NO

3 4 5 6 7 2, R = H, R = NO2, R = H, R = NO2, R = H

R

E OX = 1.33 V

R1 = H, R2 = Me, R3 = NO2, R4 = NO2, R5 = H, R6 = Me, R7 = NO2

E OX = 1.33 V

R1 = H, R2 = Br, R3 = NO2, R4 = Br, R5 = Br, R6 = NO2, R7 = H R1 = H, R2 = Cl, R3 = NO2, R4 = Cl, R5 = NO2, R6 = Cl, R7 = NO2

E OX = 1.37 V

Ar

–e–

Ar

Ar

Stilbenes

Ar

+H2O

+

–H+

E OX = 1.56 V

Ar Ar

Ar +H2O –H+

Ar

OH OH

Ar

SCHEME 15.7

Ar

+ + OH

–H+

OH +

Ar

–H+ O

Ar + O

O

Aldehydes

Proposed mechanism of an oxidative cleavage of stilbenes.

© 2016 by Taylor & Francis Group, LLC

+

Ar

Ar + OH

OH Ar

–e–

Ar

OH

Ar –e–

Ar

–e–

O

OH

537

Electron Transfer–Catalyzed Reactions R1 R2 R6

R3

R5 R4 Substrates

1.3 V vs Ag/AgNO3 4–12 F 8 (Cat.) 0.1 M LiClO4 CH3CN/H2O

R1

R4

R2

R5 + O

R3

O

R6

(+)C–Pt(–) Oxidation potentials (vs Ag/AgNO3)

Yield (%)

R1 = H, R2 = CF3, R3 = H, R4 = H, R5 = CHO, R6 = H

EOX = 1.46 V

96

R1 = H, R2 = CF3, R3 = H, R4 = H, R5 = CF3, R6 = H

EOX = 1.50 V

96

R1 = H, R2 = CN, R3 = H, R4 = H, R5 = CN, R6 = H

EOX = 1.55 V

95

EOX = 1.59 V

95

EOX = 1.70 V

93

EOX = 1.73 V

95

1 = CF

R

2 3 4 5 6 3, R = H, R = CF3, R = H, R = CHO, R = H

1 = H, R2 = CN, R3 = H, R4 = CF

R

R1 = CF

5 6 3, R = H, R = CF3

2 3 4 5 6 3, R = H, R = CF3, R = CF3, R = H, R = CF3

SCHEME 15.8 Anodic cleavage of electron-deficient stilbenes using tris(4-methyl-2-nitrophenyl)amine (8) as an oxidative mediator.

triarylamine oxidative mediators offered not only significant improvement of the current efficiencies but also effective inhibition of the side reactions. Several electron-deficient substituted stilbenes could be introduced into the electrochemical equivalent of ozonolysis that afforded a single aldehyde from symmetrical stilbenes or an equimolar mixture of two aldehydes from asymmetrical stilbenes in excellent yields. Based on the oxidation potentials, the mechanism of the indirect electrochemical cleavage of the electron-deficient substituted stilbenes was proposed (Scheme 15.9). Because the oxidation potential Anodic oxidation

Bulk oxidation O2N R1 N

R2 NO2

O2N

R6

R3

R5 R4

8 +e– –e–

–e– O2N

R5

+

R1

N O2N

NO2

R4

R6 Cleavages

R2 R3

+

Anode

SCHEME 15.9 Proposed mechanism of an anodic cleavage of electron-deficient stilbenes using tris(4methyl-2-nitrophenyl)amine (8) as an oxidative mediator.

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538

Organic Electrochemistry O

O

O

–e–

Ar S Ar S-Arylthiobenzoates

Ar

S Ar +

Ar

+ Ar S OMe

SCHEME 15.10

Ar

+MeOH –H+

Ar

S

Ar

OMe O

O –e–

+MeOH –H+

OMe

Methyl benzoates

+

S

Ar

OMe

Proposed mechanism of an oxidative methoxydesulfurization of S-arylthioarylates.

of 8 was lower than those of the stilbenes, 8 was selectively oxidized at the surface of the anodes even in the presence of the stilbenes to give the radical cation of 8, which then functioned as a bulk oxidant to generate the radical cation of the stilbenes, leading to the cleavages. After the oxidation of the stilbenes, the neutral 8 was reoxidized at the surface of the anode to catalyze the reactions. The comparison of indirect electrochemical transformations with their direct variants has also been reported using the electrochemical methoxydesulfurization of S-arylthioarylates as models [45]. The oxidation of S-arylthioarylates gave the corresponding radical cations, which were then trapped by methanol, resulting in electrochemical methoxydesulfurizations to afford the corresponding methyl arylates (Scheme 15.10). When fluoride ion was used as a nucleophile instead of methanol, electrochemical fluorodesulfurizations could also be induced [46]. The use of tris(2,4-dibromophenyl)amine (9) as an oxidative mediator produced considerable improvement in the reactions (Scheme 15.11). Because the oxidation potential of 9 was lower than those of the S-arylthioarylates, the anodic potentials during electrolysis decreased even under constant current conditions. The S-arylthioarylates were oxidized by the radical cation of 9 in the bulk electrolyte solution to drive the reactions, while neutral 9 was reoxidized through the anodic oxidation. Thus, only a catalytic amount of 9 was required to effectively complete the reactions.

C.

TEMPO

It has been well established that 2,2,6,6-tetramethylpiperidine-1-oxyl, known as TEMPO, is a typical example of a stable free radical that serves as a useful oxidant in organic synthesis and has also been studied theoretically [47–55]. In particular, TEMPO functions as an effective cocatalyst in combination with copper catalysts, enabling aerobic oxidation of a wide variety of alcohols to the corresponding carbonyl compounds [56–58]. TEMPO has also been combined with palladium catalysts under electrolytic conditions as an oxidative mediator [59–62]. Thus, TEMPO was anodically oxidized to give the corresponding N-oxoammonium cation, which could regenerate reactive cationic palladium complexes through bulk oxidation, leading to various palladium-catalyzed transformations (Scheme 15.12). In these cases, only catalytic amounts of both TEMPO and palladium catalysts were required to construct TEMPO/palladium double mediatory systems. In this context, for example, the palladium-catalyzed electrochemical cross-couplings of terminal alkynes (10) and arylboronic acids (11) were demonstrated using 4-(benzyloxy)-2,2,6,6tetramethylpiperidin-1-oxyl (12) as an oxidative mediator (Scheme 15.13) [63]. Both ethynylarenes and arylboronic acids could be introduced into the reaction to construct the corresponding diarylacetylenes.

D.

HALIDE IONS

Several halide ions can also be used as simple oxidative mediators in indirect electrochemical transformations [64,65]. Thus, halide ions, typically derived from supporting electrolytes, are oxidized at the anodes to generate the corresponding halonium ions, which then function as oxidizing reagents

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539

Electron Transfer–Catalyzed Reactions 10 mA/cm2 5F 9 (Cat.) 0.2 M Bu4NBF4 MeOH/CH2Cl2/CH3CN (+)Pt–Pt(–)

R1 S O

R2

O Yieldb (%)

62

31

R1 = H, R2 = Cl

70

40

R1 = H, R2 = OMe

80

52

1 = Me, R2 = H

R

56

37

R1 = OMe, R2 = H

67

36

b

Br

N Br

Br

R1 = H, R2 = H

a

Br

OMe

Yielda (%)

Substrates

Br

R1

Br 9

Yields of the indirect electrochemical transformations using 9. Yields of the direct electrochemical transformations without 9.

Anodic oxidation Br

Bulk oxidation Br

Br N

R1

Br

Br

S O

Br 9 –e–

R2

+e– –e– Br

Br

Br R1

+

N

+

Methoxydesulfurizations

S

Br

Br

O

R2

Br Anode

SCHEME 15.11 Anodic methoxydesulfurization of S-arylthioarylates using tris(2,4-dibromophenyl)amine (9) as an oxidative mediator.

in bulk electrolyte solutions. For example, the indirect electrochemical 2,5-dimethoxylation of furans has been chosen as a model to demonstrate the effectiveness of the use of bromide ions as oxidative mediators (Scheme 15.14) [66]. As described earlier, electron transfer between two solid phases is highly restricted; thus, solid phase–bound substrates are rarely oxidized or reduced at the surface of the electrodes in general. Meanwhile, 2,5-dimethoxylation of solid phase–supported furans (13) took place efficiently with the use of bromide ions as oxidative mediators, which could be easily combined with a multistep solid-phase synthesis [67]. The anodically generated bromonium ions functioned as oxidizing reagents in bulk solution.

E.

IODOARENES

Hypervalent iodines have been well established as easily handled oxidants that possess unique reactivities [68–70]. One of the most commonly used hypervalent iodines is 1,1,1-triacetoxy-1,1dihydro-1,2-benziodoxol-3(1H)-one, known as Dess–Martin periodinane, which can oxidize various primary alcohols into the corresponding aldehydes [71–76].

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540

Organic Electrochemistry Anodic oxidation

Bulk oxidation

N

Subtrate

1/2 Pd(0)

O TEMPO –e–

+e–

–e–

+e–

+ N

Product

1/2 Pd(II)

O Anode

SCHEME 15.12 Schematic illustration of an indirect anodic transformation using a TEMPO/palladium double mediatory system.

R2

R1 + 10

B(OH)2 11

50 mA/cm2 4F Pd(OAc)2, 12 (Cat.) DBU 0.2 M NaClO4 CH3CN/H2O (+)Ag–Pt(–)

Substrates

Yield (%)

R1 = H, R2 = tBu R1 = H, R2 = Cl

99 (42)a 91 (84)a

R1 = H, R2 = Ac

76 (77)a

R1 = Me, R2 = Me R1 = H, R2 = OPh

99 91

R1 = N2O, R2 = Ac

76

a

OBn

R1

N O 12

R2

In the absence of 12.

Anodic oxidation

Bulk oxidation OBn R2

R1 1/2 Pd(0)

N O –e–

12 OBn + N O

Anode

+ 10

+e– –e–

+e–

B(OH)2 11

R1

1/2 Pd(II) R2

SCHEME 15.13 Anodic cross-coupling of terminal alkynes (10) and arylboronic acids (11) using a 4-(benzyloxy)-2,2,6,6-tetramethylpiperidin-1-oxyl (12)/palladium double mediatory system.

© 2016 by Taylor & Francis Group, LLC

541

Electron Transfer–Catalyzed Reactions

O

15 mA/cm2 40F 0.2 M Bu4NBr MeOH/Dioxane

O 14

O

MeO

OMe

O

O

O

(+)C–C(–)

Bulk oxidation

Anodic oxidation

O

Br– O –2e–

O 14

+2e–

MeO

O

Br+

OMe O

O Anode

SCHEME 15.14 Anodic 2,5-dimethoxylation of solid phase–supported furans (14) using a bromide ion as an oxidative mediator. 5 mA/cm2 4F 15 (Cat.)

N N

S 16

R

Et3N-3HF (+)Pt–Pt(–)

l F

N N

Substrates

Yield (%)

R = COOEt R = CN

87 (31)a 72

S

N R

N+

O

(CF3SO2)2N– 15

a Yield without 15.

SCHEME 15.15 Anodic fluorination of 2-pyrimidylsulfides (16) using an ionic liquid-supported iodobenzene (15) as an oxidative mediator.

Meanwhile, several hypervalent iodoarenes can also be obtained through the anodic oxidation of the corresponding iodoarenes, which can then be utilized as oxidative mediators in indirect electrochemical transformations. In particular, the anodic oxidation of iodoarenes in the presence of fluoride ions gave the corresponding hypervalent iodoarene difluorides, which served as oxidative mediators [77]. For example, ionic liquid–supported iodobenzene (15) has been prepared as the oxidative mediator for the indirect electrochemical fluorination of 2-pyrimidylsulfides (16) (Scheme 15.15) [78]. A mechanism of the indirect electrochemical fluorination of the 2-pyrimidylsulfides was proposed based on the oxidation potentials (Scheme 15.16). The oxidation potential of 15 was lower than that of 16; thus, 15 was selectively oxidized at the surface of the anodes in the presence of fluoride ions to give the corresponding hypervalent iodoarene difluoride, which then functioned as a bulk oxidative fluorinating reagent to induce the electrochemical fluorination of 16. Through the fluorination of the 2-pyrimidylsulfides, 15 was regenerated to catalyze the reactions.

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542

Organic Electrochemistry

Anodic oxidation

l

Bulk oxidation

l N

N S 16

N 15

N+

=

O

(CF3SO2)2N–

R

15

+2e–, –2F–

–2e–, +2F–

F I

F

N F

N

S

R

Anode

SCHEME 15.16 Proposed mechanism of an anodic fluorination of 2-pyrimidylsulfides (16) using an ionic liquid-supported iodobenzene (15) as an oxidative mediator.

S

R

5 mA/cm2 4F 17

S

18

R

l F F 17

0.1 M Et4NCl Et3N-5HF (+)Pt–Pt(–) Substrates R=H R=H R=H R=H R=F R = Cl R = Me R = OMe

Yield (%) 42a 40b 51c 86 84 86 67 70

a

Yield without Et4NCl and 17. Yield without Et4NCl. c Yield without 17. b

SCHEME 15.17 Anodic fluorodesulfurization of cyclic dithioacetals (18) using a solid phase–supported iodoarene (17)/chloride ion double mediatory system.

Moreover, iodoarenes have also been effectively combined with halide ions to construct double mediatory systems. For example, chloride ions were anodically oxidized to generate chloronium ions, which then oxidized iodoarenes to give the corresponding hypervalent iodoarenes (Scheme 15.17) [79]. Chloronium ions acted as homogeneous oxidizing reagents in bulk electrolyte solution; thus, solid phase–supported iodoarenes (17) could also be oxidized to induce anodic fluorodesulfurization of cyclic dithioacetals (18). In this case, the solid phase–supported hypervalent iodoarenes could be utilized as practical oxidative fluorinating reagents. The indirect electrochemical fluorodesulfurization of 18 occurred effectively only in the presence of both chloride ions and 17. Therefore, the chloride ions/17 double mediatory system could be proposed (Scheme 15.18). The chloride ions derived from the supporting electrolyte were oxidized at the anode to play the role of initial oxidative mediator. Compound 17 then served as a secondary oxidative mediator to direct the reactions.

© 2016 by Taylor & Francis Group, LLC

543

Electron Transfer–Catalyzed Reactions Anodic oxidation Bulk oxidation S

l

S

Cl– 17 –

–

–2e

–

–

R

18

R

–2e–, –2F–

+2e –2e , +2F

F l

+

Cl

F F F

Anode

SCHEME 15.18 Proposed mechanism of an anodic fluorodesulfurization of cyclic dithioacetals (18) using a solid phase–supported iodoarenes (17)/chloride ion double mediatory system.

F.

O -AMINOPHENOLS

Creating biomimetic systems in indirect electrochemical transformations has proven to be an intriguing challenge. It is well known that benzoquinone structures play critical roles in biological electron transfer processes. The benzoquinones are reduced to give the corresponding hydroquinones, which can be reoxidized to benzoquinones to function as redox mediators in biological systems. Based on this fact, several o-iminoquinones generated through the anodic oxidation of the corresponding o-aminophenols have been utilized as oxidative mediators in indirect electrochemical transformations to realize exquisite reactions [80–83]. For example, the anodic oxidation of 1-(3-amino-2,4-dihydroxyphenyl)ethanone (19) gave the corresponding o-iminoquinone, which induced chemoselective imine formation, leading to the N-alkylation of various primary amines through the cathodic reduction (Scheme 15.19) [84]. In this reaction, 19 functioned as an effective biomimetic oxidative mediator (Scheme 15.20). The anodically generated o-iminoquinone, which catalyzed the formation of the corresponding imine of the substituted benzylamine, was quantitatively re-reduced to 19 after the reaction, driving the reaction cycle. Thus, only a catalytic amount of 19 was required to consume the starting substituted benzylamine.

NH2

NH2

+

0.1 M LiClO4 MeOH (+)Pt–Pt(–)

R

–1.6 V vs SCE 1h 0.1 M LiClO4 MeOH (+)Pt–Hg(–)

OH O

0.6 V vs SCE 3h 19 (Cat.)

H2N N R

HO 19

N H R

Substrates R=F R = Me R = OMe

Yield (%) 68 62 65

SCHEME 15.19 Anodic N-alkylation of primary amines using 1-(3-amino-2,4-dihydroxyphenyl)ethanone (19) as a biomimetic oxidative mediator.

© 2016 by Taylor & Francis Group, LLC

544

Organic Electrochemistry Anodic oxidation

Bulk oxidation O

OH H2N

NH2

+

NH2

R

HO 19 –2e–, –2H+

+2e–, +2H+ OH

O N

HN R O Anode

SCHEME 15.20 Proposed mechanism of an anodic N-alkylation of primary amines using 1-(3-amino-2,4dihydroxyphenyl)ethanone (19) as a biomimetic oxidative mediator.

III.

REDUCTIVE MEDIATORS

A.

GENERAL MECHANISM

In contrast to oxidative mediators, reductive mediators are activated through cathodic reduction and the resulting intermediates function as subsequent reductants (Scheme 15.21). While several oxidative mediators have been established to expand the possibility of anodic oxidation processes, reductive mediators have been more limited. In this section, indirect electrochemical transformations using reductive mediators, including nickel(II) salens and carboranes, are discussed with their use in recent synthetic examples.

B.

NICKEL(II) SALENS

Cathodic reductive cyclizations have been studied extensively to build a wide variety of ring systems in unique ways [85–87]. Cathodic reduction of electron-deficient olefins tethered to electrophiles, such as aldehydes or ketones, gave the corresponding intramolecular cyclized products through carbon–carbon bond formations. Such reductive cyclizations could also be induced using oneelectron reductants, including tributyltin hydrides and samarium diiodides. Meanwhile, the use of nickel(II) salens as reductive mediators has been extensively studied both synthetically and mechanistically [88–91]. For example, the cathodic reduction of electron-deficient olefins in the presence of Cathodic reduction

Bulk reduction Mediator (oxidized)

+e–

Substrate (oxidized) –e–

Mediator (reduced)

+e– Substrate (reduced)

Chemical transformations

Cathode

SCHEME 15.21

Schematic illustration of indirect cathodic transformation using reductive mediators.

© 2016 by Taylor & Francis Group, LLC

545

Electron Transfer–Catalyzed Reactions –2.10 V vs Ag/AgNO3 2F 20 (Cat.)

O R2

R1

0.1 M Bu4NBr CH3CN (+)Pt–RVC(–)

Substrates

N

OH R1

N Ni

O

R2

O 20

Yield (%)

R1 = Me, R2 = COOMe

70

R1 = H, R2 = CN

73

Bulk reduction

Cathodic reduction N

N

O

Ni R

O

O 20

+e–

1

R2

–e– •–

N

N Ni O

O

OH R1 R2

Cathode

SCHEME 15.22 Cathodic intramolecular cyclization of electron-deficient olefins using nickel(II) salen (20) as a reductive mediator.

nickel(II) salen (20) gave the corresponding intramolecular cyclized products even under a constant potential condition at significantly lower potentials than the reduction potentials of the substrates (Scheme 15.22) [92].

C.

O -CARBORANES

Carboranes are polyhedral clusters composed of boron and carbon atoms. In particular, icosahedral carboranes [93,94], known as o-carboranes, have high thermal and chemical stabilities, and thus have been studied intensively in many fields of chemistry, including medicinal [95,96], material [97–99], and physical [100–103]. The redox properties of the o-carboranes have also been found to produce both stable radical anions and dianions, which enable their use as reductive mediators in indirect electrochemical transformations. Although polyaromatic hydrocarbons, such as fullerenes [104] and transition metal ions, can also be used as reductive mediators, their low solubilities in typical organic solvents limit their use in indirect electrochemical transformations. Meanwhile, several substituents can be introduced into the two carbon atoms of o-carboranes to realize facile modulations of their solubilities and reduction potentials. For example, it has been demonstrated that 1,2-diphenyl-o-carborane (21) could serve as an efficient reductive mediator in the indirect electrochemical debromination of 1,2-dibromo-1,2-diphenylethane (22) (Scheme 15.23) [105]. Based on their reduction potentials, the dianion of 21 acted as a bulk reductant to induce the debromination. After the reduction of 22, the radical anion of 21 was reduced again at the surface of the cathodes to catalyze the reactions. In contrast, no direct electrochemical debromination of 22 took place at this potential.

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Organic Electrochemistry –1.3 V vs SCE 2F 21 (Cat.) 0.1 M Et4NClO4 DMF (+)Pt–C(–)

Br Br 22

Ph

Ph 21 ( = BH)

Amount of 21 (mol%)

Yield (%)

0 20 10 1

0 65 93 95

Bulk reduction

Cathodic reduction Ph

•–

Br Br

Ph

22

+e–

–e– Ph

2–

Ph Cathode

SCHEME 15.23 Cathodic debromination of 1,2-dibromo-1,2-diphenylethane (22) using 1,2-diphenyl-ocarborane (21) as a reductive mediator.

IV. OXIDATIVE CHAIN REACTIONS A.

GENERAL MECHANISM

As indicated earlier, electrochemical chain reactions can be defined as autoactivating chemical transformations induced by direct heterogeneous electron transfer. Generally, there are three stages in the chain reactions: initiation, propagation, and termination. For example, in a radical chain reaction, the initiation stage consists of the use of chemical radical initiators to generate the reactive radical substrates. The resulting reactive radical substrates then induce autoactivating chemical transformations, which is the propagation stage. Finally, in the termination stage, the resulting radicals are consumed by several chemical reactions such as radical coupling. In the case of electrochemical chain reactions, initially, the starting substrate is activated through electron transfer at the surface of the electrode to give the reactive intermediate, which is then responsible for the subsequent activation in the homogeneous bulk electrolyte solution to drive the chain cycle without the use of additional electricity. In analogy with the redox mediators, there are also two types of electrochemical chain reactions: oxidative and reductive. Oxidative chain reactions are initiated through anodic oxidation and the resulting intermediates function as subsequent oxidants, which are regenerated through the reactions to drive the chain cycle (Scheme 15.24). In this section, typical oxidative chain reactions directed by reactive radical cations and cations are discussed in detail using inter- and intramolecular carbon–carbon bond formations as models.

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Electron Transfer–Catalyzed Reactions Bulk oxidation reaction

Anodic oxidation Starting substrate

Chemical transformations

–

–e

Reactive intermediate

Reactive intermediate

Anode

SCHEME 15.24

Schematic illustration of an anodic chain reaction.

B. RADICAL CATION CHAIN REACTION Anodic oxidation of electron-rich olefins has proven to be an effective approach for reversing the polarity of olefins and triggering radical cation-based intramolecular cyclizations, which provide powerful tools to construct various ring systems [106–110]. In particular, electrochemical olefin cross-couplings between electron-rich olefins and unactivated olefin nucleophiles took place through radical cation chain reactions involving intermolecular carbon–carbon bond formations [111–114]. For example, the anodic oxidation of 3,4-dihydro-2H-pyran (23) in the presence of several substituted allylbenzenes (24) affords the corresponding cycloadducts (Scheme 15.25) [115]. In these cases, the electron densities of the aromatic rings were critical for the reactions. Based on their oxidation potentials, an oxidative chain reaction mechanism was proposed (Scheme 15.26). Because the oxidation potential of 23 was lower than that of 24, 23 was first selectively oxidized at the anode to generate the corresponding radical cation. The radical cation of 23 was then trapped by 24 to form a relatively long-lived aromatic radical cation, which was finally reduced by the starting 23 to construct the corresponding cycloadduct with regeneration of the radical cation of 23, enabling the radical cation chain mechanism.

C.

CATION CHAIN REACTION

As described earlier, radical chain reactions have been widely utilized in both academic and industrial synthetic chemistry to efficiently produce various organic molecules. In contrast, R2

R1 + R3 24

O 23

1.0 V vs Ag/AgCl 0.5 F 1.0 M LiClO4 CH3NO2 (+)C–C(–)

Substrates R 1 = H , R 2 = H, R3 = H R

1= H ,

R 2 = Me, R3 = H

R1

R2

R3

O

Yield (%) 0 13

R 1 = Me, R 2 = Me, R3 = H

60

R 1 = Me, R 2 = Me, R3 = Me

84

R 1 = H, R 2 = OMe, R3 = H

94

SCHEME 15.25 Anodic intermolecular cyclization of 3,4-dihydro-2H-pyran (23) and substituted allylbenzenes (24).

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Organic Electrochemistry Anodic oxidation

Radical cation chain reaction R2 O 23

O 23

R1 O

R3 –e– +e–

–e–

R1

R2 + +

+

O

O

O

R3

Anode R2

R1

R3 24

SCHEME 15.26 Proposed mechanism of an anodic intermolecular cyclization of 3,4-dihydro-2H-pyran (23) and substituted allylbenzenes (24).

F

O R

O

25 (Cat.)

ArS 26

0.1 M Bu4NB(C6F5)4 CH2Cl2

R

Substrates

Yield (%)

R = c-Hex R = C7H15

88

R = CH2Ph

75

R = Ph

65

R = CH2CH2OMe

62

+ S S S

SAr

F

= ''ArS+'' F 25

82

SCHEME 15.27 Anodic intramolecular cyclization of thioacetals (26).

cation chain reactions are still limited. In organic electrochemistry, it has been well established that several carbocations could be accumulated through the anodic oxidation of substrates in the absence of nucleophiles at low temperature, known as the “cation pool” method [116–118]. In particular, ArS(ArSSAr) +, an equivalent of ArS+, could be generated through the anodic oxidation of ArSSAr in dichloromethane electrolyte solution [119–122]. For example, “ArS+” (25) was accumulated at a low temperature to induce the intramolecular carbon–carbon bond formations of various thioacetals (26), which occurred through cation chain reactions (Scheme 15.27) [123]. Initially, 25 was accumulated through the anodic oxidation of the corresponding ArSSAr without 26 in dichloromethane electrolyte solution at a low temperature (Scheme 15.28). After the electrolysis, 26 was then added into the solution to give the corresponding alkoxycarbenium ions and ArSSAr. The alkoxycarbenium ions were trapped by intramolecular olefin nucleophiles to form the corresponding reactive carbocation intermediates, which finally reacted with ArSSAr to regenerate 25 with the formation of the cyclization products, realizing the cation chain mechanism.

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Electron Transfer–Catalyzed Reactions Cation chain reaction Anodic oxidation O R

ArSSAr

ArSSAr

O

ArS 26

ArSSAr

R

SAr

–2e–, H+ ''ArS+'' 25

+O

''ArS+'' 25

O R

R

+

''ArS+'' 25

Anode

SCHEME 15.28

Proposed mechanism of an anodic intramolecular cyclization of thioacetals (26).

V. REDUCTIVE CHAIN REACTIONS A. GENERAL MECHANISM While oxidative chain reactions are initiated by anodic oxidations, reductive chain reactions are triggered through cathodic reductions. The cathodic reductions generate reactive intermediates that function as subsequent reductants, which are regenerated through the reactions to drive the chain cycle (Scheme 15.29). Representative reductive chain reactions are aromatic nucleophilic substitutions, known as SRN1 reactions (Scheme 15.30) [124,125]. Initially, the haloarene was reduced to the corresponding radical anion, from which a halide ion was subsequently eliminated to form the aryl radical. The radical was then trapped by a nucleophile to give the corresponding radical anion, which was finally oxidized by the starting haloarene to construct the corresponding substituted arene with regeneration Anodic oxidation

Bulk reductive reaction Starting substrate

Chemical transformations

+e–

Reactive intermediate

Reactive intermediate Anode

SCHEME 15.29

Schematic illustration of a cathodic chain reaction. Ar X

Ar X

Ar Nu

+e–

+e– –

–

–X–

SCHEME 15.30

Ar X = Haloarene

–e– –

Ar

+Nu–

Proposed mechanism of a reductive aromatic nucleophilic substitution of haloarenes.

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Organic Electrochemistry

of the radical anion of the haloarene, enabling the SRN1 mechanism. In the termination stage, the aryl radical was consumed through several chemical reactions such as trapping by solvents. For example, the cross-couplings of aryl Grignard reagents and haloarenes have been reported to take place through SRN1 pathways [126]. In this section, reductive chain reactions are discussed in detail using electrochemical SRN1 reactions as models.

B.

ELECTROCHEMICAL SRN1 REACTION

Several types of SRN1 reactions can be induced through electrochemical approaches. The cathodic reduction of haloarenes is responsible for the initiation of these reactions [127,128]. Generally, the use of reductive mediators is effective for the initiation of electrochemical SRN1 reactions. For example, the cross-couplings of 1-iodo-2-(trifluoromethyl)benzene (28) and the substituted 1H-imidazoles (29) in the presence of phthalonitrile (30) as a reductive mediator occurred through SRN1 pathways (Scheme 15.31) [129]. The electrochemical SRN1 reactions were initiated through the cathodic reduction of 30 (Scheme 15.32). The cathodically generated radical anion of 30 functioned as a bulk reductant to reduce 28, forming the corresponding radical anion. An iodide ion was then eliminated from the radical anion to give the corresponding radical, which was then trapped by the anion of 29 as a nucleophile. The cross-coupling was completed through the oxidation of the resulting radical anion by the starting 28 with regeneration of the radical anion of 28, realizing the reductive chain reaction. N

l + CF3 28

R N H 29

–1.8 V vs SCE 0.8F 30 (Cat.) 0.1 M Et4NBF4 DMSO (+)Pt–C(–)

Substrates

CN

NH N

R CN 30

CF3

Yield (%)

R = CHO R = C6H4OMe

42 55

SCHEME 15.31 Cathodic cross-coupling of 1-iodo-2-(trifluoromethyl)benzene (28) and substituted 1H-imidazoles (29) using phthalonitrile (30) as a reductive mediator. Cathodic reduction

Bulk reduction CN

l

Reductive chain reaction l

CN

CF3

CF3

NH R

30 +e–

28

N

28

CF3

–e– +e–

+e– –e–

CN –

l

CN

CF3

–

l

–

NH

–

R N

CF3

CF3

Cathode –l–

CF3

N N

R

SCHEME 15.32 Proposed mechanism of a cathodic cross-coupling of 1-iodo-2-(trifluoromethyl)benzene (28) and substituted 1H-imidazoles (29) using phthalonitrile (30) as a reductive mediator.

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Electron Transfer–Catalyzed Reactions

VI. A.

MEASURINg REDOX POTENTIALS GENERAL ASPECTS

As discussed in this chapter, not only redox potentials of substrates but also those of redox mediators are of considerable significance in the design of indirect electrochemical transformations. Redox potential can be defined as a measure of the tendency of compounds to gain or lose an electron and thereby be reduced or oxidized. Generally, the selectivity of the electrochemical transformations is determined by the value of redox potentials, for example, the substrate of lower oxidation potential can be oxidized preferentially even in the presence of the excess amount of the substrate of higher oxidation potential. In this section, direct and indirect techniques to measure redox potentials are discussed in detail.

B.

DIRECT MEASURING

Cyclic voltammetry is the most basic and widely used technique to acquire information about electron transfer of the compounds directly, for example, their redox potentials. Although the value of redox potentials is highly dependent upon the measuring conditions, including electrode materials and supporting electrolytes, the redox potentials of both ions and radical ions have been systematically reviewed in the literatures. Cyclic voltammetry also provides information about reaction mechanisms of electrochemical transformations [130–132]. For example, the electrochemical transformations of α-tocopherol (31) were studied in detail by using cyclic voltammetry (Scheme 15.33) [133].

C.

INDIRECT MEASURING

While the redox potentials of both ions and radical ions have been systematically reviewed in the literatures, those of radicals are almost unknown because their reactivities are extremely high. On the other hand, the Marcus theory has been well established to relate the kinetics of the electron transfer to its thermodynamics, which are determined by the redox potentials. Accordingly, the latter values of highly reactive radicals can be estimated reasonably through the investigation of the reaction kinetics. In this context, the “competition” method has been developed to indirectly determine the redox potentials of radicals [134–137]. For example, alkyl radicals can be generated through the indirect cathodic reduction of the corresponding haloalkanes using anthracene (32) as a reductive mediator (Scheme 15.34) [138]. Initially, 32 was reduced to the corresponding radical anion, which then reduced the haloalkanes to their radical anions. The corresponding alkyl radicals were formed through the subsequent elimination of halide ions. In electrochemical SRN1 reactions, +

OH

OH

–e– R

O

+e–

+e–

O

R

2+

OH

–e– O

R

31 +H+ –H+

+H+ –H+

O– R

O

+H+ –H+

O

–e– +e– R

O

+e–

R = (CH2CH2CH2CHCH3)2CH3

SCHEME 15.33

Redox reaction of α-tocopherol (31).

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O

–e–

+ R

O

552

Organic Electrochemistry R

33 Bulk reduction

Cathodic reduction

+H+ –

R R–Br 32 –e– +e–

+e– –

k1 –Br–

–

R +

k2 Cathode R– + 32 +H+

R–H + 32

SCHEME 15.34

Redox reaction of haloalkanes using anthracene (32) as a reductive mediator.

the generated radical was trapped by a nucleophile. On the other hand, in this case, the radical anion of 32 would be the only reaction partner for the radicals. Thus, there were two main different pathways for the reaction between the alkyl radicals and the radical anion of 32, including crosscoupling and reduction. While the corresponding anion of the alkyl anthracene (33) was formed through the cross-coupling, the reduction gave the anion of alkane and 32. Because the rate constant of the cross-coupling k1 could be measured experimentally, the rate constant of the reduction k2 could also be calculated, enabling the reasonable estimation of the redox potentials of the alkyl radicals based on the Marcus theory.

VII. CONCLUSIONS As described in this chapter, indirect electrochemical transformations using redox mediators can significantly expand the possibilities of organic electrochemistry. To date, a wide variety of both oxidative and reductive mediators are available to synthetic chemists to realize useful electrochemical transformations in unique ways. Homogeneous electron transfer between substrates and activated redox mediators in bulk electrolyte solution appears to present a thermodynamic paradox in that the substrates can be effectively oxidized or reduced even at lower potentials than those required to be oxidized or reduced through heterogeneous electron transfer at the surface of electrodes. Additionally, both oxidative and reductive chain reactions can be induced through the electrochemically generated reactive intermediates in indirect ways, which is intriguing both synthetically and mechanistically. Therefore, indirect electrochemical transformations should find further potential applications to enable green sustainable reactions in both academic and industrial fields of chemistry.

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Organic Electrochemistry Wu, X.; Davis, A. P.; Lambert, P. C.; Steffen, L. K.; Toy, O.; Fry, A. J. Tetrahedron 2009, 65, 2408–2414. Wu, X.; Davis, A. P.; Fry, A. J. Org. Lett. 2007, 9, 5633–5636. Halas, S. M.; Okyne, K.; Fry, A. J. Electrochim. Acta 2003, 48, 1837–1844. Shen, Y.; Hattori, H.; Ding, K.; Atobe, M.; Fuchigami, T. Electrochim. Acta 2006, 51, 2819–2824. Shen, Y.; Suzuki, K.; Atobe, M.; Fuchigami, T. J. Electroanal. Chem. 2003, 540, 189–194. Jeena, V.; Robinson, R. S. Chem. Commun. 2012, 48, 299–301. Hoover, J. M.; Stahl, S. S. J. Am. Chem. Soc. 2011, 133, 16901–16910. Han, B.; Wang, C.; Han, R.-F.; Yu, W.; Duan, X.-Y.; Fang, R.; Yang, X.-L. Chem. Commun. 2011, 47, 7818–7820. Ishii, K.; Kubo, K.; Sakurada, T.; Komori, K.; Sakai, Y. Chem. Commun. 2011, 47, 4932–4934. Hirota, M.; Furihata, K.; Saito, T.; Kawada, T.; Isogai, A. Angew. Chem. Int. Ed. 2010, 49, 7670–7672. Kusamoto, T.; Kume, S.; Nishihara, H. Angew. Chem. Int. Ed. 2010, 49, 529–531. Pouliot, M.; Renaud, P.; Schenk, K.; Studer, A.; Vogler, T. Angew. Chem. Int. Ed. 2010, 49, 6037–6040. Bardelang, D.; Banaszak, K.; Karoui, H.; Rockenbauer, A.; Waite, M.; Udachin, K.; Ripmeester, J. A.; Ratcliffe, C. I.; Ouari, O.; Tordo, P. J. Am. Chem. Soc. 2009, 131, 5402–5404. Kusamoto, T.; Kume, S.; Nishihara, H. J. Am. Chem. Soc. 2008, 130, 13844–13845. Figiel, P. J.; Sibaouih, A.; Ahmad, J. U.; Nieger, M.; Räisänen, M. T.; Leskelä, M.; Repo, T. Adv. Synth. Catal. 2009, 351, 2625–2632. Mannam, S.; Alamsetti, S. K.; Sekar, G. Adv. Synth. Catal. 2007, 349, 2253–2258. Jiang, N.; Ragauskas, A. J. J. Org. Chem. 2006, 71, 7087–7090. Mitsudo, K.; Shiraga, T.; Kagen, D.; Shi, D.; Becker, J. Y. Tanaka, H. Tetrahedron 2009, 65, 8384–8388. Mitsudo, K.; Shiraga, T.; Tanaka, H. Tetrahedron Lett. 2008, 49, 6593–6595. Mitsudo, K.; Kaide, T.; Nakamoto, E; Yoshida, K.; Tanaka, H. J. Am. Chem. Soc. 2007, 129, 2246–2247. Mitsudo, K.; Kumagai, H.; Takabatake, F.; Kubota, J.; Tanaka, H. Tetrahedron Lett. 2007, 48, 8994–8997. Mitsudo, K.; Shiraga, T.; Mizukawa, J.; Suga, S.; Tanaka, H. Chem. Commun. 2010, 46, 9256–9258. Tajima, T.; Imai, N.; Nakajima, A.; Kurihara, H.; Fuchigami, T. J. Electroanal. Chem. 2006, 593, 43–46. Baba, D.; Fuchigami, T. Electrochim. Acta 2003, 48, 755–760. Nad, S.; Breinbauer, R. Angew. Chem. Int. Ed. 2004, 43, 2297–2299. Nad, S.; Roller, S.; Haag, R.; Breinbauer, R. Org. Lett. 2006, 8, 403–406. Zhdankin, V. V.; Stang, P. J. Chem. Rev. 2008, 108, 5299–5358. Zhdankin, V. V.; Stang, P. J. Chem. Rev. 2002, 102, 2523–2584. Stang, P. J.; Zhdankin, V. V. Chem. Rev. 1996, 96, 1123–1178. Nicolaou, K. C.; Sugita, K.; Baran, P. S.; Zhong, Y.-L. J. Am. Chem. Soc. 2002, 124, 2221–2232. Nicolaou, K. C.; Baran, P. S.; Zhong, Y.-L.; Sugita, K. J. Am. Chem. Soc. 2002, 124, 2212–2220. Barrett, A. G. M.; Hamprecht, D.; Ohkubo, M. J. Org. Chem. 1997, 62, 9376–9378. De Munari, S.; Frigerio, M.; Santagostino, M. J. Org. Chem. 1996, 61, 9272–9279. Dess, D. B.; Martin, J. C. J. Am. Chem. Soc. 1991, 113, 7277–7287. Dess, D. B.; Martin, J. C. J. Org. Chem. 1983, 48, 4155–4156. Fuchigami, T.; Fujita, T. J. Org. Chem. 1994, 59, 7190–7192. Sawamura, T.; Kuribayashi, S.; Inagi, S.; Fuchigami, T. Org. Lett. 2010, 12, 644–646. Sawamura, T.; Kuribayashi, S.; Inagi, S.; Fuchigami, T. Adv. Synth. Catal. 2010, 352, 2757–2760. Largeron, M.; Chiaroni, A.; Fleury, M.-B. Chem. Eur. J. 2008, 14, 996–1003. Xu, D.; Chiaroni, A.; Fleury, M.-B.; Largeron, M. J. Org. Chem. 2006, 71, 6374–6381. Largeron, M.; Neudorffer, A.; Fleury, M.-B. Angew. Chem. Int. Ed. 2003, 42, 1026–1029. Largeron, M.; Neudorffer, A.; Vuilhorgne, M.; Blattes, E.; Fleury, M.-B. Angew. Chem. Int. Ed. 2002, 41, 824–827. Largeron, M.; Fleury, M.-B. Org. Lett. 2009, 11, 883–886. Little, R. D. Chem. Rev. 1996, 96, 93–114. Fry, A. J.; Little, R. D.; Leonetti, J. J. Org. Chem. 1994, 59, 5017–5026. Little, R. D.; Fox, D. P.; Van Hijfte, L.; Dannecker, R.; Sowell, G.; Wolin, R. L.; Moens, L.; Baizer, M. M. J. Org. Chem. 1988, 53, 2287–2294. Dunach, E.; Franco, D.; Olivero, S. Eur. J. Org. Chem. 2003, 1605–1622. Esteves, A. P.; Freitas, A. M.; Medeiros, M. J.; Pletcher, D. J. Electroanal. Chem. 2001, 499, 95–102. Fielder, S. S.; Osborne, M. C.; Lever, A. B. P.; Pietro, W. J. Am. Chem. Soc. 1995, 117, 6990–6993. Isse, A. A.; Gennaro, A.; Vianello, E. Electrochim. Acta 1992, 37, 113–118. Miranda, J. A.; Wade, C. J.; Little, R. D. J. Org. Chem. 2005, 70, 8017–8026. Plesek, J. Chem. Rev. 1992, 92, 269–278. Bregadze, V. I. Chem. Rev. 1992, 92, 209–223.

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Section IV Organic Electrochemical Reaction Types

© 2016 by Taylor & Francis Group, LLC

16

Cleavages and Deprotections Ole Hammerich

CONTENTS I. II.

III.

IV.

V. VI.

Introduction ....................................................................................................................... 561 Carbon–Carbon Bonds ..................................................................................................... 562 A. Reductive C–C Cleavages ......................................................................................... 562 1. Esters .................................................................................................................. 562 2. Nitriles ................................................................................................................ 562 3. Ketones ............................................................................................................... 563 4. 1,2-Diols .............................................................................................................564 5. Halogen Compounds ..........................................................................................564 6. Hydrocarbons .....................................................................................................564 B. Oxidative C–C Cleavages ......................................................................................... 565 1. Carboxylic Acids ................................................................................................ 565 2. Ketones ............................................................................................................... 565 3. 1,2-Diols and Related Hydroxy Compounds ...................................................... 565 4. Alkoxy Compounds............................................................................................ 566 5. Alkenes ............................................................................................................... 567 6. Hydrocarbon Single Bonds ................................................................................ 568 Carbon–Nitrogen Bonds ................................................................................................... 569 A. Reductive C–N Cleavages ......................................................................................... 569 1. Quaternary Ammonium Ions ............................................................................. 569 2. Carboxamides and Related Compounds ............................................................ 570 3. Aminoketones .................................................................................................... 571 4. Amines ............................................................................................................... 571 5. Azides ................................................................................................................. 572 6. Diazo Compounds .............................................................................................. 572 7. Nitro Compounds ............................................................................................... 573 B. Oxidative C–N Cleavages ......................................................................................... 573 1. Amides ............................................................................................................... 573 2. Amines ............................................................................................................... 574 Carbon–Phosphorus Bonds ............................................................................................... 574 A. Reductive C–P Cleavages.......................................................................................... 574 1. Phosphonium Ions .............................................................................................. 575 2. Phosphonates ...................................................................................................... 575 3. Phosphine Oxides ............................................................................................... 575 4. Phosphines.......................................................................................................... 575 Carbon–Arsenic Bonds ..................................................................................................... 576 A. Reductive C–As Cleavages ....................................................................................... 576 Carbon–Oxygen Bonds ..................................................................................................... 577 A. Reductive C–O Cleavages ......................................................................................... 577 1. Acylphosphonium Salts ...................................................................................... 577 2. Carboxylates and Carbonates ............................................................................. 577

559

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560

VII.

VIII. IX. X.

XI. XII. XIII.

XIV.

Organic Electrochemistry

3. Sulfonates and Related Esters ............................................................................ 579 4. β-Arylthio- and β-Alkylthioesters and -Alcohols .............................................. 579 5. β-Nitroesters ....................................................................................................... 580 6. Phosphates .......................................................................................................... 580 7. Hydroxyketones .................................................................................................. 581 8. Hydroxysulfones ................................................................................................. 581 9. Alcohols.............................................................................................................. 581 10. Acyclic Ethers and Glycosides ........................................................................... 581 11. 1,3-Dioxolanes and 1,3-Dioxanes ....................................................................... 583 12. Oxiranes ............................................................................................................. 583 B. Oxidative C–O Cleavages ......................................................................................... 584 1. Enol Esters and Related Compounds ................................................................. 584 2. Ethers.................................................................................................................. 584 Carbon–Sulfur Bonds ....................................................................................................... 585 A. Reductive C–S Cleavages.......................................................................................... 585 1. Sulfonium Ions ................................................................................................... 585 2. Thiol Esters ........................................................................................................ 587 3. Thioethers (Sulfides) .......................................................................................... 588 4. Sulfoxides ........................................................................................................... 589 5. Sulfones .............................................................................................................. 591 6. Thiocyanates ...................................................................................................... 594 B. Oxidative C–S Cleavages .......................................................................................... 594 1. Thiol Esters ........................................................................................................ 594 2. Thioethers (Sulfides) .......................................................................................... 595 3. Disulfides............................................................................................................ 596 Carbon–Selenium and Carbon–Tellurium Bonds ............................................................. 596 A. Reductive C–Se and C–Te Cleavages........................................................................ 596 B. Oxidative C–Se and C–Te Cleavages ........................................................................ 597 Carbon–Halogen Bonds .................................................................................................... 597 A. Reductive C–Hal Cleavages ...................................................................................... 597 Nitrogen–Nitrogen Bonds ................................................................................................. 597 A. Reductive N–N Cleavages ......................................................................................... 598 1. Carboxylic Acid Hydrazides and Azides ........................................................... 598 2. Azines and Related Compounds ........................................................................ 598 3. Diazo Compounds .............................................................................................. 598 B. Oxidative N–N Cleavages ......................................................................................... 599 Nitrogen–Phosphorus Bonds............................................................................................. 599 A. Reductive N–P Cleavages ......................................................................................... 599 Nitrogen–Oxygen Bonds ................................................................................................... 599 A. Reductive N–O Cleavages......................................................................................... 599 Nitrogen–Sulfur Bonds .....................................................................................................600 A. Reductive N–S Cleavages .........................................................................................600 1. Sulfonamides ......................................................................................................600 2. Sulfenamides ......................................................................................................602 3. Sulfimides (Sulfilimines) ....................................................................................602 B. Oxidative N–S Cleavages..........................................................................................602 1. Sulfinamides .......................................................................................................602 2. Sulfenamides ......................................................................................................603 Nitrogen–Halogen Bonds ..................................................................................................603 A. Reductive N–Hal Cleavages ......................................................................................603

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561

XV.

Phosphorus–Phosphorus Bonds ........................................................................................603 A. Reductive P–P Cleavages ..........................................................................................603 XVI. Phosphorus–Oxygen Bonds ..............................................................................................603 A. Reductive P–O Cleavages .........................................................................................603 B. Oxidative P–O Cleavages..........................................................................................604 XVII. Phosphorus–Sulfur Bonds ................................................................................................604 A. Reductive P–S Cleavages ..........................................................................................604 XVIII. Phosphorus–Selenium Bonds............................................................................................605 A. Reductive P–Se Cleavages ........................................................................................605 XIX. Phosphorus–Halogen Bonds .............................................................................................605 A. Reductive P–Hal Cleavages ......................................................................................605 XX. Arsenic–Arsenic Bonds ....................................................................................................605 A. Reductive As–As Cleavage .......................................................................................605 XXI. Oxygen–Oxygen Bonds ....................................................................................................606 A. Reductive O–O Cleavages ........................................................................................606 XXII. Oxygen–Sulfur Bonds .......................................................................................................606 A. Reductive O–S Cleavages .........................................................................................606 1. Alkoxysulfonium Ions ........................................................................................606 2. Sulfonates ...........................................................................................................606 3. Sulfinates ............................................................................................................607 4. Sulfenates ...........................................................................................................607 5. Sulfoxides ...........................................................................................................608 XXIII. Sulfur–Sulfur Bonds .........................................................................................................608 A. Reductive S–S Cleavages ..........................................................................................608 B. Oxidative S–S Cleavages ..........................................................................................608 XXIV. Sulfur–Halogen Bonds ......................................................................................................609 A. Reductive S–Hal Cleavages ......................................................................................609 1. Sulfonylhalides ...................................................................................................609 2. Sulfenylchlorides ................................................................................................609 XXV. Selenium–Selenium and Tellurium–Tellurium Bonds ...................................................... 610 A. Reductive Se–Se and Te–Te Cleavages ..................................................................... 610 B. Oxidative Se–Se and Te–Te Cleavages ..................................................................... 610 References ...................................................................................................................................... 610

I. INTRODUCTION Most electrochemical conversions include the cleavage of a covalent bond as a part of the route from starting materials to products, for instance, substitution reactions, deprotonation of radical cations, and elimination of anionic leaving groups during many reductions just to mention a few examples. Still, the cleavage step is often the characteristic feature of the electrochemical process and for this reason chapters dedicated to cleavages are often included in electrochemistry books [1,2], and it has been decided by the editors of this book to maintain this tradition also in this fifth edition of Organic Electrochemistry. Owing to the widespread occurrence of electrochemical cleavages, some tough decisions had to be made when preparing this chapter in order to limit the size of the presentation. Only the cleavage of bonds between what traditionally is being considered as nonmetals is included. Cleavages observed during the electrochemistry of organometallic compounds are treated separately in Chapter 36. Even so, some of the cleavages observed for organic compounds containing only nonmetallic heteroatoms belong more naturally in other chapters in this book; this concerns the reductive cleavage of the carbon–halogen bond that is treated in detail in Chapters 24 and 25 as well

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Organic Electrochemistry

as organoboron and organosilicon compounds, and to some extent organophosphorus compounds, that are included in the chapter dedicated to the electrochemistry of organoelemental compounds (Chapter 35). In addition to the chapters already mentioned, the reader may also find discussions of cleavage reactions including the group 16 elements in Chapter 27 and N–N bond in Chapters 29 and 30. In many cases, cleavage processes have been studied in great detail and the results have been discussed within the theory of dissociative and stepwise electron transfers. Such aspects are treated in Chapters 13 and 14.

II. CARbON–CARbON bONDS A.

REDUCTIVE C–C CLEAVAGES

1. Esters Reduction of tetramethyl ethane-1,1,2,2-tetracarboxylate, in a DMF–10% water mixture at a mercury cathode at 60–65°C, proceeds with cleavage of the central C–C bond to dimethyl malonate Equation 16.1 [3]. Related substituted and cyclic compounds react similarly, the latter accompanied by ring opening. The yields are generally in the range 90–100%. O

O

O

O

O

O

+2e–, +2H+

O

(16.1)

2

O O

O

O

O

Analogous to these is the retro-Bingel reaction observed for the dianions of bis(methoxycarbonyl) methanofullerenes and phosphorylated methanofullerenes Equation 16.2 [4,5]. 2– ROOC

P(O)(OR)2 P(O)(OR)2

2–

+2e–, +2HB

+ CH2

–2B–

(16.2)

COOR

2. Nitriles Reductive cleavage of the C–CN bond is observed under nonacidic conditions during the reduction of aromatic nitriles in which the aromatic ring is π-electron deficient owing to the presence of a heteroatom, such as in cyanopyridines [6,7], or to an electron-withdrawing substituent, such as in benzenedicarbonitriles [8,9] and alkoxycarbonylbenzonitriles [10–12]. Reduction of 2- and 4-cyanopyridine in nonacidic aqueous solution at the potential of the second CV peak gives pyridine and cyanide ion according to the stoichiometry shown in Equation 16.3. In contrast, reduction in acidic solution proceeds to the aminomethylpyridines [6,7]; the radical anion of isophthalonitrile undergoes a practically irreversible dimerization in aprotic solvents [11]. CN +2e–, +H2O

+ CN–

–OH– N

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N

(16.3)

563

Cleavages and Deprotections

The reduction of aliphatic nitriles at a zinc cathode in DMF/Et4NOTs leads to decyanation in yields that often exceed 70% [13]. The reaction in the last step is a convenient procedure for α-alkylation of amines, Equation 16.4, the first step being an anodic α-methoxylation (see Chapter 19). –2e–, –2H+, MeOH

Me3SiCN, TiCl4 OMe

N

N

COOR΄

COOR΄

– CN

(16.4) +2e

–, +H+

–CN–

R N

COOR΄

CN

COOR΄

R–X –X–

Base N

N

N

CN

COOR΄

R

COOR΄

3. Ketones Reduction of α-diketones [14] in DMF, in the presence of oxygen, results in cleavage of the central C–C bond and, after alkylation, in the formation of the corresponding esters (62–98%), Equation 16.5. O

O

O O 1) +2e–, +O2

(16.5)

2) Mel O O

The reaction is suggested to include the attack of electrogenerated superoxide ion on the carbonyl group followed by a series of steps that include cleavage of the central C–C bond. Symmetrical α-diimines are reductively cleaved to the corresponding amides (40–70%) in a similar fashion [15]. The reaction between electrogenerated superoxide and chalcones leads to a series of cleavages and eventually the formation of the corresponding carboxylic acids (Scheme 16.1) [16]. The radical anion of 2-phenylbenzoylcyclopropane undergoes irreversible cleavage with the formation, after further reduction, of 1,4-diphenyl-l-butanone (66%), Equation 16.6 [17]. In contrast, the cleavage observed for the radical anion of the unsubstituted benzoylcyclopropane is reversible; the ring-opened radical anion reacts with the substrate to give dimeric or trimeric products. O–

O Ph

Ph

O – + Ph +e , +2H

Ph

Ph Ph

(16.6)

The radical anion of phenyltritylketone, prepared by reduction in DMF, undergoes cleavage of the C–C bond according to Equation 16.7 at the time scale of a preparative electrolysis [18]. Protonation of the trityl anion leads to triphenylmethane (80–85%), whereas the arylcarbonyl radical reacts

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564

Organic Electrochemistry O–

O O2 Ar΄

Ar

Ar΄ O

O2 Ar

SCHEME 16.1

Ar΄

COO– + Ar

Ar

CHO

O

ArCH2COO– + ArCOO–

CHO

Superoxide induced cleavage reactions.

with DMF resulting in the formation of methylene dibenzoate as the other major product (~85%). 2-Naphthyltritylketone reacts similarly. Ph

Ph

Ph Ph

O

slow

Ph

O

+

(16.7)

Ph Ph

Ph

4. 1,2-Diols Pinacols are reduced in protic media to the corresponding hydrocarbons in a process that involves cleavage of both C–O and C–C bonds [19,20]. An example, the reduction of 9,9′,10,10′-tetrahydro9,9′-bianthracene-9,9′-diol to 9,10-dihydroanthracene (90%), is shown in Equation 16.8 [19]. HO 2.8 F

(16.8)

2

2

In strongly basic media, the pinacol is cleaved by base to the radical anion of the parent ketone in a nonelectrochemical process. 5. Halogen Compounds Substituted diphenylamines may be prepared by reduction of N-(4-methyl-4-(trichloromethyl) cyclohexa-2,5-dienylidene)anilines in a process that includes the reductive two-step elimination of the CCl3 group, the first being a C–Cl cleavage and the second an elimination of dichlorocarbene, Equation 16.9 [21]. Me NH2 + O

X

Me

p-TolSO3H

X

N

CCl3

+2e– –Cl–

Me X

N CCl2

CCl3

+H+ – :CCl2

X

H N

Me

(16.9) 6. Hydrocarbons Reduction of 9,9′-bianthryl in DMF leads to anthracene [22]; in this case via cleavage at the dianion stage and subsequent protonation of the strongly basic anthracene anion by solvent/supporting electrolyte components; the radical anion of 9,9′-bianthryl is a stable species in aprotic solvents.

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Cleavages and Deprotections

B. OXIDATIVE C–C CLEAVAGES 1. Carboxylic Acids The electrochemical oxidative cleavage of aliphatic carboxylic acids to alkyl radicals, R·, and carbon dioxide is a classic reaction in organic electrochemistry. The radical may either dimerize to R–R (the Kolbe reaction) or be further oxidized to a carbocation, R+, that subsequently reacts with a suitable nucleophile (see Chapter 33 for details). Here only two examples will be given. Electrolysis of a vicinal dicarboxylic acid may be used to introduce a double bond in a ring system as, for example, in the transformation of a 3-oxo-2-azabicyclo[2.2.2]oct-5-ene-7,8-dicarboxylic acid to a 2-azabicyclo[2.2.2]octa-5,7-dien-3-one, Equation 16.10 [23]. O

O

Me

Me

N

N

–2e–, –2H+, –2CO2

COOH

(16.10)

COOH

Instead, the conversion of l-nonyl-3-oxocyclohexanecarboxylic acid ethylene ketal to methyl 5-nonyl-5-hexenoate by oxidation at a carbon anode in methanol containing K2CO3 includes a carbocation intermediate, Equation 16.11 [24]. COOH + C9H19

–

–2e , –H , –CO2

O

O

C9H19

+

O

O

(16.11) C9H19

C9H19 MeOH + O

COOCH3 O

2. Ketones Ketones with branching in the α-position may be cleaved at a platinum anode in MeCN– LiC1O4 [25]. A mechanism that involved generation of a carbenium ion by α-cleavage of the ketone radical cation, Equation 16.12, analogous to reactions in photochemistry and mass spectroscopy, was proposed. O –2e– R3C R3C

Me NHCOMe + Me

R3C+ + Me

+

C

O

MeCN, H2O

(16.12)

COOH

3. 1,2-Diols and Related Hydroxy Compounds Vicinal glycols and related monohydroxy compounds may be cleaved to carbonyl compounds or acetals, Equation 16.13, by oxidation at a carbon anode in MeOH/Et4NOTs [26,27]. The anodic oxidation does not show the stereochemical limitations usually observed when the cleavage is

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Organic Electrochemistry

carried out with chemical oxidizing agents. 1,2-Diols may also be cleaved by indirect oxidation by using the periodate–iodate redox system as a mediator [28,29]. R3

R2

R3

R2

R

OR

RO

(16.13)

O + O

R4

R1

1

4

R

Symmetrical ketones may be prepared in this way, for example, by alkylation of methyl methoxyacetate with RMgX to the dialkyl methoxymethylcarbinol that is then oxidized [27]. Yields are typically around 80%. Unsymmetrical ketones may be obtained from symmetrical ketones by anodic oxidation of the enol ethers to the α-methoxyketone [27,30], reaction with R′MgX and then anodic oxidation of the glycol monomethyl ether. Alternatively, the ketone may be oxidatively aminated, alkylated at C═O, and oxidized, Equation 16.14, Y═OCH3 or NR2. O

OH R΄MgX

RCH2

RCH2

CHR

O ~3F, MeOH

CHR R΄

RCH2

+ RCOOH

(16.14)

R΄

Y

Y

Unexpectedly, it was observed that for the oxidation of the racemic and meso forms of 4,4′-dihydroxyhydrobenzoin at an anode of mild steel at pH 13.5 only the racemic form was cleaved to p-hydroxybenzaldehyde at a measurable rate, whereas the meso form could be recovered unchanged [31]. An indirect oxidative cleavage by oxidation with “Cl+” has been reported for 2-amino-1cycloalkanols, resulting in the formation of the corresponding keto nitriles [32]. Related to this is the oxidative cleavage of N-o-phenylbenzoyl prolinols in MeOH to the α-methoxylated pyrrolidine derivative [33] and of the retro Paterno–Büchi reaction observed during the oxidative cleavage of oxetanes mimicking DNA (6–4) photoproducts [34]. β-Hydroxyhydroxylamines may be cleaved to aldehydes and oximes upon anodic oxidation at pH 8, Equation 16.15 [35]. R΄

R΄ –2e–, –2H+

R

NHOH

CH

R pH 8

OH

(16.15)

NOH

CHO + R˝

R˝

4. Alkoxy Compounds 1,1,2-Trimethoxycyclohexane may be oxidized to the acetal ester [36]. The starting material for this reaction may itself be prepared by anodic addition of MeOH to cyclohexanone enol ether leading to the overall reaction shown in Equation 16.16. OMe OMe

MeO

O

MeO MeOH –2e–, –2H+

H2O

MeO OMe

–2e–, –2H+ H

(16.16)

OMe

Related to this reaction is the oxidation of veratrole in MeOH/KOH to, among other products, 5,5,6,6-tetramethoxy-l,3-cyclohexadiene and hexamethyl cis,cis-orthomuconate,

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Cleavages and Deprotections

Equation  16.17 [37]; carried out independently as a separate step, the latter cleavage gives a yield of 77%. It is noteworthy that only the least stable cis,cis isomer is formed; this may indicate that the intermediates are short-lived and/or adsorbed on the electrode during the fixation of the geometry of the product. OMe

MeO OMe OMe

OMe MeOH

OMe

MeOH

–2e–, –2H+

OMe

–2e–, –2H+

OMe OMe

MeO

(16.17)

OMe OMe

Alkoxycyclopropane rings may be opened by anodic oxidation, which of the bonds that are cleaved depends on the substituents and the experimental conditions. Thus, 7,7-dichloro-1ethoxybicyclo[4.1.0]heptane [38] gives methyl 7,7-dichlorohept-6-enoate (86%) by anodic oxidation at a carbon anode in MeOH in the presence of lutidine, Equation 16.18, whereas 7,7-dichloro-ltrimethylsilyloxybicyclo[4.1.0]heptane upon oxidation in MeOH at –13 to –10°C in the presence of Fe(NO3)3 [39] gives methyl 2-oxocyclohexanecarboxylate (72%), Equation 16.19. Cl

Cl

–2e–, –2H+ MeOH/lutidine

Cl

(16.18) COOMe

Cl

OEt

Cl Cl

COOMe i = 18 mA

(16.19)

MeOH O

OSiMe3

Anodic oxidation of l-methoxy-2-phenylthiocyclopropanes [40] in MeOH/K2CO3 proceeds with cleavage of the ring between C–l and C–2; the product is the corresponding 2-[methoxy(phenylthio) methyl]alkanone, 79%, Equation 16.20. The oxidation of 2-phenylthio-1-cycloalkanols proceeds in a fashion similar to that of diketones [41]. OMe

SPh

2.3F SPh

(16.20)

MeOH O OMe

5. Alkenes Alkenylarenes may undergo cleavage of the double bond by oxidation at a dimensionally stable anode or a Pb anode in MeOH or EtOH containing CF3COONa as the electrolyte [42]. The reaction proceeds through an initial oxidative addition of RO – to the double bond followed by cleavage to the dialkyl acetal in a reaction similar to that in Equation 16.13. The cleavage of a 1,2-diol is an intermediate step in the oxidation of stilbenes carrying electron-withdrawing substituents to the corresponding aldehydes [43]. The oxidative addition of –OH and the subsequent cleavage were carried out in MeCN/water using triarylamines as redox catalysts.

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Organic Electrochemistry

A rather complicated reaction sequence, including the oxidative cleavage of a C–C bond carrying a methoxy substituent, has been proposed to account for conversion of 1,2-,5,6-bis[trimethylene] cyclooctatetraene to methyl 5-(2,3-dihydro-1H-inden-5-yl)-5-oxopentanoate by oxidation in MeOH, Equation 16.21 [44]. OMe

1) Electrocyclization 2) –e–

2 MeOH –2e–, –2H+ OMe OMe MeO

MeOH –e–, –2H+ OMe

OMe

+

+

(16.21)

+

OMe

OMe

MeO

+

MeO

OMe

OMe

MeOH –2e–, –H+ O

O OMe

Oxidative cleavage of alkenes and cycloalkenes to carboxylic acids and diacids, isolated as the methyl esters, in 70–90% yields has been accomplished by using a IO 4 −/RuCl3 mediator system [45]. For example, cyclohexene could be converted in this way to dimethyl hexanedioate in 76% yield. Oxiranes may be ring opened in acidic methanol to diols that in turn may be cleaved anodically; this sequence has been used to convert a propylene side chain, via electrochemical epoxidation (not shown below), to an aldehyde, Equation 16.22 [46]. O

O

O THF, 10%H2SO4

O

Ph

O

O

Ph

3–7F MeOH, H2SO4

O

Ph

(16.22)

CHO

HO OH

6. Hydrocarbon Single bonds Cleavage of the central C–C bond has been observed during the oxidation of 1,2-di-(p-tert-butylphenyl)ethane using boron-doped diamond electrodes [47]. The reaction products, p-tert-butylbenzaldehyde dimethyl acetal and p-tert-butylbenzyl methyl ether were believed to arise via a mechanism including attack by methoxy radicals resulting from the oxidation of MeOH at the BDD electrodes.

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Cleavages and Deprotections

III. CARbON–NITROgEN bONDS A. REDUCTIVE C–N CLEAVAGES 1. Quaternary Ammonium Ions Salts of tetraalkylammonium, R4N+, and related ions are widely used as supporting electrolytes in organic electrochemistry in aprotic solvents (see Chapter 7). One reason for that is that R4N+ is difficult to reduce and this allows for electrochemical studies of a broad range of substrates without interference from the reduction of the supporting electrolyte cation. However, at sufficiently low potentials, R4N+ is reduced and the nature of the reduction products turns out to depend on both the cathode material and the concentration of residual water in the solvent [48–52]. At mercury, tetraalkyl ammonium amalgams are formed [48,53], whereas reductive cleavage of the type shown in Scheme 16.2 for R = Et is observed at most other cathode materials [49,50,52,54]. The reaction includes a dissociative electron transfer as the first step followed by reduction of the ethyl radical. The resulting ethane anion reacts with Et4N+ in a Hofmann elimination producing ethene and ethane. Also, the cathodic breakdown of ionic liquid cations may involve C–N cleavage [55]. Early work [56] has shown that anilinium salts with three N-alkyl groups are reduced at a Pb cathode to benzene and the corresponding tertiary amine. Reduction of N-allyl-N,N-dimethylanilinium ions (or N-benzyl-N,N-dimethylanilinium ions) gives propylene (or toluene) and dimethylaniline [57]; in general, it is found that ease of cleavage follows the order benzyl > allyl > phenyl > alkyl. The same type of cleavage is observed in DMF and it was found that quaternary ammonium ions containing allyl, benzyl, fluorenyl, acenaphthenyl, benzhydryl, and cinnamyl groups gave propylene, toluene, fluorene, acenapthene, diphenylmethane, and a mixture of allylbenzene and propenylbenzene, respectively [58]. Occasionally, dimers such as bibenzyl derived from the neutral free radical formed in the initial electron transfer process have been observed [59–61]. The formation of a mixture of meso- and d,l-2,3-diphenylbutane in the reduction of enantiomerically pure d-α-phenylethyl-l-trimethylammonium nitrate in DMF [59] as well as data from ESR spectroscopy [61] support a mechanism including radical intermediates. Reductive ring opening of anilinium salts in which one of the N-alkyl groups is attached to the 2-position of the benzene ring is a useful synthetic procedure, Equation 16.23 [62]. Other heterocyclic quaternary ammonium salts behave similarly [63]. (CH2)n

n = 2, 90%

+2e–, +H+

(CH2)n

+

n = 4, 90%

N

N+

+e– N

+e– H3C

CH2

H3C

+ H3C

CH2

CH2– H

N+

+ H3C

CH–2

N

Reductive cleavage of tetraethylammonium ion.

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H C

+ H

SCHEME 16.2

(16.23)

n = 3, 88%

N

+ H3C

C H

CH3

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Organic Electrochemistry

2. Carboxamides and Related Compounds The products resulting from reduction of amides depends on the acidity of medium [6]. For example, reduction of isonicotinic anilide gives 4-pyridinylmethanol at pH 5, whereas a mixture of N-(4-pyridinylmethyl)aniline and 4-pyridinylmethanol is obtained in 1 N HCl. A gem-amino alcohol that may lose either the amine forming an aldehyde, or water forming an imine, is a likely intermediate. The aldehyde is then reduced to the alcohol, whereas the imine is reduced to an amine, Equation 16.24. A similar gem-amino alcohol is formed during the reduction of oxaziridines [64]. OH CONHPh

CHNHPh H

N+

+

CH2NH2–Ph + H2O

+2e–, +3H+

+2e–, +2H+ N+

N+

H

H

(16.24) H

N+

CH2OH + Ph

NH3+

Simple carboxamides are difficult to reduce [65] and for that reason not suitable in protection/ deprotection chemistry of amines or amino acids. However, the corresponding nitro derivatives are easily cleaved, in acidic media [66] to the deprotected amine and benzisoxazolone via the hydroxylamine, Equation 16.25, or in DMF [67]. O

O NHR

NO2

+4e–, +4H+ –H2O

O NHR

NHOH

O + RNH2

(16.25)

N H

Benzyloxycarbonyl derivatives of primary and secondary amines may be reduced at vitreous carbon cathodes in DMF with cleavage to toluene and the free amine in good yields (>80%), Equation 16.26, [68]. By using a high-surface area, Pd cathode cleavage may be carried out under mild conditions; selectivities in this type of deprotection are excellent [69]: O C6H5CH2O

+2e–, +2H+ NHR

C6H5CH3 + CO2 + RNH2

(16.26)

The reduction may be made easier by using 4-nitrobenzyloxycarbonyl [70] or cinnamyloxycarbonyl [71] derivatives making these derivatives suitable in amine protection/deprotection chemistry. Another approach includes the reductive deprotection of allyl carbamates in a [Ni(bipy)3](BF4)2 catalyzed decarboxylation-type process [72]; the reactions were carried out in DMF in a singlecompartment cell with a consumable Zn rod anode. tert-Butoxycarbonyl-substituted acyl amides may be selectively deacylated by electrochemical reduction [73]. Chloro-, bromo-, and iodoethoxycarbonyl as well as trifluoro-, trichloro, and tribromoethoxycarbonyl have been used as easily removed protection groups in peptide synthesis [74,75]. N-acylureas undergo reductive cleavage of a C–N bond to the corresponding isocyanate and an amide anion. Hydrolysis and protonation gave the final products (Scheme 16.3) [76]. N-(Diethylaminothiocarbamoyl)benzamidines undergo C–N cleavage upon cathodic reduction. It was found that the nature of the substituent at the amidine nitrogen affected which of the C–N bonds that undergoes cleavage as illustrated by Scheme 16.4 [77].

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571

Cleavages and Deprotections O NH

R

C

O +e–

N

R

R

NH

N

C

R

–

–½H2 COR΄

COR΄ O

N

R

C

Li+ N–

O + R

R΄

C

H2O

H2O

O RNH2 + CO2

SCHEME 16.3

H N

R

C

R΄

Reductive cleavage of N-acylureas.

NH +HS

R = H, n-C3H7

CH2

N(Et)2

N H N N H

N(Et)2 S

R

+4e–, +4H+

CH2

R

N

+ R

SCHEME 16.4

N(Et)2

H2N

H

R = Ph

S

Reductive cleavage of N-(diethylaminothiocarbamoyl)benzamidines.

3. Aminoketones The reduction of α-aminocarbonyl compounds, Equation 16.27 [78–80], or α-aminonitriles [81] proceeds similarly to the reductive cleavage of the α-hydroxy compounds, in this case with elimination of the amine. R

R˝

R

R˝

+2e–, +2H+ NR΄2 O

R΄˝

H + R΄2NH O

(16.27)

R΄˝

An exception to this general scheme was found for the reduction of two alkaloids, quininone and dihydrocinchoninone, in 35% sulfuric acid; under these conditions, reduction took place at the carbonyl group leaving the C–N bonds intact [82]. Reduction in MeCN resulted in “normal” C–N cleavage. 4. Amines Reduction of 4-aminobenzonitrile in DMF resulted in the formation of the radical anion of 4,  4′-dicyanobiphenyl [8]. This product was proposed to arise from elimination of NH2− at the radical anion stage followed by dimerization of the neutral radical and further reduction.

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Organic Electrochemistry

2,6-Di-tert-butyl-4-phenylcyclohexa-2,5-dienones may be used as protecting groups for amino acids and peptides [83]; protection is easily accomplished by anodic oxidation of 3,5-di-tertbutylbiphenyl-4-ol in the presence of the amino acid in CH2Cl2/Et4NBF4 [84] and deprotection by reductive cleavage in MeOH/HCl [83], Equation 16.28. OH

O

MeOH, HCl

R΄ Ph

NH

R΄

Ph and/or OH

–1.1 V vs Ag/Ag+

COOR

+

NH2

COOR

(16.28)

R˝ Cl

R˝

Ph

5. Azides Reduction of azides activated by a neighboring >C=O or >C=N–NH–CS–NH 2 function leads to cleavage of the C–N bond and elimination of azide ion, Equation 16.29. The final products are imidazole derivatives (70–80%) resulting from a complex cascade of follow-up reactions [85]. O O Ar

O

+2e– CH2N3

–N3–

Ar

Ar

CH2N3

N

Many steps Ar

CH2–

N

COAr

(16.29)

H

6. Diazo Compounds The reactivity of the radical anions resulting from electrochemical reduction of organic diazo compounds is often sufficiently low to allow for studies of the kinetics and mechanisms of the follow-up reactions by CV using moderate sweep rates. In a number of cases such as azibenzil [86,87] and diethyl diazomalonate [87], it has been shown that the radical anion decomposes in a unimolecular reaction generating the corresponding carbene radical anions and dinitrogen, Equation 16.30. Hydrogen-bonding to water deliberately added to the MeCN solution facilitates the cleavage reaction and so does the neighboring carbonyl group. –

– O

O

+ N2

(16.30)

N2

The radical anion of 2,3,4,5-tetraphenyldiazocyclopentadiene was found to decompose in a similar fashion [88]. A more complicated behavior was observed for the decomposition of the

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Cleavages and Deprotections

diazodiphenylmethane radical anion to largely diphenylmethane and benzophenone azine [89], a reaction that had earlier been suggested to involve unimolecular decomposition as in Equation 16.30 [90]. However, although the radical anion decay was observed to give rise to a first order rate law, a rate decrease of a factor of 20 was found when using CD3CN instead of CH3CN as the solvent indicating the participation of hydrogen-abstraction from the solvent as an important part of the decay mechanism that includes several parallel paths. In contrast, the 9-diazofluorene radical anion undergoes rate determining dimerization as the first step in a complex chain reaction that leads to the product, fluorenone azine [91]. Reduction of the related bis(diazo) compounds of indenofluorenes results in the formation of oligomeric polyazines with chain lengths varying from 3–4 up to around 16 [92]. Related to these reactions is the reductive conversion of diethyl diazo(phenyl)methylphosphonate to diethyl benzylphosphonate in aqueous dioxane, Equation 16.31 [93]. OEt

OEt O

O P

P

+2e–, +2H+

OEt

OEt

+ N2

(16.31)

N2

7. Nitro Compounds It has been observed in a number of cases that the reduction of organic nitro compounds is accompanied with cleavage of the C–N bond and the formation of nitrite ions. A few examples have been reported for the reduction of aromatic nitro compounds, including 1,4-dinitrobenzene [94] and l,2,4,5-tetrafluoro-3,5-dinitrobenzene [95]. However, more often this type of cleavage is met during the reduction of compounds in which the nitro group is attached to an aliphatic carbon as in nitrocumenes [96,97] and 2-(4-nitrophenyl)-2-nitropropane [98] and in the reduction of purely aliphatic compounds such as 2,2-dinitropropane [99,100] and 1,1-dinitrocyclohexane [101–104]. The reader is referred to Chapter 30 for details.

B. OXIDATIVE C–N CLEAVAGES 1. Amides Electrochemical oxidation of peptides and proteins has been shown to lead specifically to C–N cleavage next to tyrosine and to tryptophan [105], here illustrated by cleavage next to tyrosine, Equation 16.32. OH

O

O

O

+

O

–2e–, –H+

O

O

O N R˝

N R΄

O

H H

N R΄

N

H H

O

+

R΄

N N H H

R˝

H2O –H+

O + H2N R˝

(16.32)

O

R˝ R΄

N H

O

Coupled to EC-MS analysis this approach was suggested as an alternative to chemical and enzymatic cleavages.

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N-(4-Methoxyphenyl)-2-azetidinones and N-(2,4-dimethoxyphenyl)-2-azetidinones may be oxidatively deprotected under mild conditions in a MeCN/H2O (1/1) mixture containing 1% LiClO4; it was suggested that C–N cleavage took place at the dication stage, Equation 16.33 [106]. R΄

R

R΄

R

R΄

R H2O

–2e– N

–H+

N

N

+

O

+

O

O

OMe OMe

OMe

OH

R΄

R H2O

+

–H+

+

O

O + MeOH

(16.33)

N O

H

2. Amines Anodic oxidations that proceed beyond the formation of radical cations and dications are usually accompanied by the liberation of protons in one or more of the steps on the way to the final products. During the investigation of the electrochemical oxidation of, for instance, aliphatic amines that are often the strongest bases in the solution, this inevitably gives rise to rather complicated mechanism schemes including acid–base reactions that involve the substrate. In addition, the oxidation of aliphatic amines is markedly dependent on the concentration of residual water owing to the possibility that reaction intermediates containing imine-like structures undergo in situ hydrolysis. As an example, tripropylamine has been found to be oxidatively cleaved in MeCN that has not been scrupulously dried according to Scheme 16.5 [107]; only the most important steps on the way to dipropylamine and propanal are shown and it is not specified either how the protons are distributed among the various amine species. N-(4-Methoxyphenyl) amines are easily oxidized owing to the presence of the methoxyphenyl group; this has been put to use in the protection/deprotection chemistry of amines [108]. When the oxidative cleavage is carried out under acidic, nonaqueous conditions, the methoxyphenyl group is converted to p-quinone. The yields obtained by electrochemical oxidation are generally higher than those obtained by cerium ammonium nitrate oxidation, typically 70–90%.

IV. CARbON–PHOSPHORUS bONDS A.

REDUCTIVE C–P CLEAVAGES

The reduction of organophosphorus compounds often includes cleavage of a C–P bond. A number of such reactions are included in a recent review [109]. The oxidation of organophosphorus compounds leads to the formation of products that do not include C–P cleavage [109,110]; accordingly, this section focuses on reductions only. –e–

(C3H7)3N (C3H7)3N

+

(C3H7)2NCHC2H5 +

(C3H7)2N = CHC2H5

SCHEME 16.5

(C3H7)3N

(C3H7)2NCHC2H5 + H+ –e–

+

(C3H7)2N = CHC2H5

H2O

Oxidative cleavage of tripropylamine.

© 2016 by Taylor & Francis Group, LLC

+

(C3H7)2NH + C2H5CHO + H+

575

Cleavages and Deprotections

1. Phosphonium Ions Phosphonium ions are more easily reduced than the corresponding ammonium ions; the reduction proceeds in a fashion similar to that of the ammonium ions and includes the cleavage of a C–P bond [49,111–116]. The effect of the electrode material, the potential, and the temperature on the relative rate of cleavage of a variety of alkyl and aryl groups has been studied in detail and it was found, for instance, that the ease of cleavage at an Hg cathode increases in the order methyl < phenyl < ethyl < n-butyl < iso-propyl < tert-butyl < benzyl [111]. The facile cleavage of the bond to the benzyl group is the basis of a convenient procedure for the preparation of mixed tertiary phosphines and quaternary phosphonium salts [111], Equation 16.34. This method also provides a simple means for preparing phosphonium salts with four different ligands that after resolution into their optical antipodes may be cleaved reductively providing optically active tertiary phosphines [117]: (PhCH2)3P

R1X –X

R2X –X–

–

(PhCH2)3R1P+

(PhCH2)2R1R2P+

+2e–, +H+ PhCH2R1R2R3P+

–PhCH3

+2e–, +H+ –PhCH3

R1R2R3P

+2e–, +H+ –PhCH3

(PhCH2)2R1P

PhCH2R1R2P

R4X –X–

R3X –X–

(16.34)

R1R2R3R4P+

Reduction of alkyltriphenylphosphonium salts in aprotic media leads to the formation of ylids [118,119], resulting from deprotonation of the starting materials by an electrogenerated base [120], in competition with the cleavage reaction [121]. Competition between ylid formation and cleavage has been observed also during the reduction of benzyl-, allyl-, cinnamyl-, and polyenylphosphonium ions [122] and of 1,2-vinylene and 1,4-butadienylene bis-phosphonium ions [123] (See Sections XVI and XIX). 2. Phosphonates Reduction of dialkyl aroylphosphonates in MeCN in the presence of benzoic acid as a proton donor gives the corresponding dialkyl α-hydroxyarylmethylphosphonates resulting from reduction of the carbonyl group [124]. This is in contrast to the reduction in the absence of acid; in this case cleavage of the C–P bond was observed resulting in the formation of the benzoins and the dialkyl hydrogenphosphonates as the only products. 3. Phosphine Oxides The reduction of triphenylphosphine oxide directly to triphenylphosphine is accompanied by a number of side reactions [125,126] and is as such not a synthetically useful reaction; in aprotic media a C–P bond is cleaved in preference to the P–O bond. (See Sections XVI and XIX.) 4. Phosphines The electrochemical reduction of Ph3P and related compounds in aprotic solvents (MeCN, DMF, and HMPA) leads to the formation of the radical anion as the first step [125–128]. The follow-up reaction is complicated, and in the presence of an R4N+ salt as the supporting electrolyte, a catalytic reduction of the R4N+ cation takes place resulting in the formation of RPh2P among other products [126] (Scheme 16.6). In addition to the steps shown, the alkyl radical may dimerize or undergo disproportionation to the corresponding alkane and alkene. The formation of biphenyl, presumably resulting from dimerization of the phenyl radicals, was reported in earlier work carried out in DMF [125]:

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Organic Electrochemistry +e– Ph3P

Ph3P –

–e–

Ph3P – + R4N+

RPh3P + R3N RPh2P + Ph

Ph3P + R Ph Ph3P – +

HB –B–

PhH

R–

HB –B–

RH

Ph3P + R

SCHEME 16.6

Ph–

Reduction of triphenylphosphine in the presence of tetraalkylammonium ions.

The reductive cleavage of a C–P bond has been observed also for the carborane shown in Equation 16.35 [129] (the noncarbon atoms in the icosahedron are all boron).

Ph2P Ph2P

H

C +2e–, +2H+

C

C H

C

–2 PPh2

(16.35)

V. CARbON–ARSENIC bONDS A.

REDUCTIVE C–AS CLEAVAGES

Quaternary arsonium ions are easier to reduce than the corresponding phosphonium ions by approximately 0.3 V [113]. The ease of cleavage of the C–As bond is mostly similar to that for the corresponding phosphonium ions and the preparative aspects are similar as well [113,130]. Tertiary arsines behave during electrochemical reduction essentially as the corresponding phosphines (see Section IV.A.3). For instance, 1-naphthyldiphenylarsine is reduced to the radical anion that subsequently undergoes cleavage of the Ph–As bond, Equation 16.36 [127]:

Ph2AsNaph

+e– Ph2AsNaph –e–

–

–Ph

PhAsNaph–

(16.36)

Cleavage of the Carom–As bond is observed during the reduction of methyl 4-(diethylarsino) benzoate, the corresponding oxide, methyl 4-(diethylarsoryl)benzoate, and the sulfide, methyl 4-(diethylarsorothioyl)benzoate. The radical anion of the resulting methyl benzoate was detected by ESR spectroscopy [131].

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Cleavages and Deprotections

VI. CARbON–OXygEN bONDS A. REDUCTIVE C–O CLEAVAGES The reductive cleavage of a C–O bond requires as the rule activation of the bond; this can be obtained in different ways. Conjugated carbinols, like allylic and benzylic alcohols, π-electrondeficient heterocyclic carbinols, and α-hydroxyketones and the ethers of these compounds, may be reductively cleaved. Unactivated aliphatic and aromatic hydroxyl groups may be reductively removed after esterification with a suitable inorganic acid whereby the hydroxyl group is transformed into a better leaving group. Reductive elimination of two vicinal hydroxyl groups or derivatives thereof may also be possible. Many of the systems described later may be used as protecting groups for alcohol functions [132,133], and the choice between them will depend, among other factors, on the presence of electrophores in the protected molecule. 1. Acylphosphonium Salts Acylphosphonium salts, obtained by constant current electrolysis of Ph3P in the presence of a carboxylic acid in CH2C12 [134], may be reduced to triphenylphosphine oxide and the aldehyde in a process that includes cleavage of the O–CO bond, Equation 16.37 [135]: +2e–, +H+

+

Ph3P

(16.37)

Ph3P = O + RCHO

OCOR

By using Ph3PH+, ClO4− as the supporting electrolyte, the reaction has been used for the conversion of α-amino acids to α-amino aldehydes [135]. Paired electrolysis in MeCN containing R3P and Et4NBr may be used in a similar way to deoxygenate primary and secondary alcohols to alkanes, Equation 16.38 [110]. The anodic process includes the formation of R3PBr2 that on reaction with R′–OH gives R′O–PR3+,Br− and elimination of R3P=O gives R′–Br that is reduced at the cathode to R′–H. R΄

OH

Paired electrosynthesis

(16.38)

R΄ H 48–96%

R3P, (R = Ph, Bu, PhO), Et4NBr

2. Carboxylates and Carbonates Alkyl and benzyl esters of aromatic [136,137] or heteroaromatic acids [138] undergo C–O cleavage upon reduction in an aprotic solvent to the carboxylate ion and alkyl or benzyl radical, here illustrated by the reduction of benzyl benzoate, Equation 16.39 [136]. The latter is further reduced to the anion under the conditions and finally protonated to toluene. O Ph

O

+2e–, +H2O O

Ph

–OH–

Ph

O–

+ H3C

Ph

(16.39)

A study of the relative rates of cleavage of a series of p-methoxycarbonylbenzyl carboxylates demonstrated that there is no simple relation between the leaving group ability and structure [139]. If a proton donor is not added, the base inevitably generated during the reduction may cause cleavage of the ester to benzoate and benzyl alcohol. The cleavage may be carried out indirectly by reduction with electrogenerated superoxide ion [140]. In that case, the reaction proceeds “only” to the carboxylate and the alcohol/alcoholate. Similarly, α-benzoyloxyacetophenone is reduced to benzoic acid and acetophenone [141] and benzyl ethyl oxalate to 2-ethoxy-2-oxoacetate and toluene [142]. During the reduction of benzyl carbonates, the intermediate benzyl anion may be trapped by carbon dioxide; the overall reaction provides an efficient synthesis of arylacetic acids in good yields [143].

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Organic Electrochemistry

The reductive cleavage of phenyl benzoate in DMF follows a different route including the initial dimerization of the radical anion followed by elimination of phenoxide ion. The resulting benzil is further reduced to the dianion, Equation 16.40 [144]. O– O 2 Ph

–

Ph O

OPh

Ph

– OPh –2 O

–O

O

O

Ph

O–

+2e–

(16.40)

Ph Ph

Ph

Ph

Ph

O–

Reductive cleavage including an initial dimerization step is observed in MeCN and DMF also for nitrobenzoates [145]. Benzil also results from the reductive cleavage of O-acylated benzaldehyde cyanohydrins [146]. The ease of reduction of 2-haloethyl esters depends on the number and type of the halogen atoms in the manner intuitively expected. Reductive cleavage of a trichloroderivative in MeOH/LiClO4 proceeds according to Equation 16.41 [74]. This type of protecting group has been employed during a nucleotide synthesis using the triester method (see Section VI.A.6). O

O +2e– –Cl–

OCH2CCl2–

Ph

OCH2CCl3

Ph

MeOH -MeO–

(16.41) O

O + Ph

87%

H2C

CCl2

+ Ph

O–

6%

OCH2CHCl2

Allyl carboxylates may be reductively cleaved to alkenes in a reaction catalyzed by Pd(0)(PPh3)4, Equation 16.42, or to the corresponding allyltrimethylsilanes if the reduction is carried out in the presence of trimethylsilyl chloride [147]. The reaction has been successfully used in protection/ deprotection chemistry as illustrated by the conversion of allyl benzoate to benzoic acid in 97% yield. R R

OAc PdLn

+2e– –AcO–

R

+ E+ –

R

E

(16.42)

AcO

In addition to allyl, the cinnamyl [148], p-tolyl [149], benzhydryl [150], triphenylmethyl [150], and 4-picolyl [151] groups are suitable in protection/deprotection chemistry as they are easily removed electrochemically. In uncatalyzed reductions, cinnamyl esters are reduced in preference to allyl esters, and cinnamyloxy carbonates are reduced easier than cinnamyloxy carbamates [148]. Propargyl acetate has been reductively cleaved to the corresponding allenes in DMF in the presence of a catalytic amount of PdCl2(PPh3)2/PPh3 in a “less than 2F” process indicating that PPh3 acts as a reducing agent through formation of an acetyltriphenylphosphonium salt that is then hydrolyzed to triphenylphosphine oxide [152]. Cleavage including an initial dimerization step is observed for the reduction in MeCN and DMF of benzoate diesters of 1,2-diols to benzil in a reaction that is likely to proceed via the formation of a cyclic 2,3-diphenyl-1,4-dioxane-2,3-diol dianion [153]. Reduction of aliphatic esters of 1,2-diols often proceeds as a reductive elimination to the corresponding alkene. For example, vicinal dioxalates are cleaved electrochemically to alkenes in DMF, Equation 16.43 [154]. The oxalates of 1,2-diols may be formed by base-catalyzed transesterification

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Cleavages and Deprotections

with diethyl oxalate [155]. Related to this is the formation of alkenes and alkanes from pinacols by reductive cleavage of cyclic carbonates and aryl boronates [156]. O– O R΄

O R΄

O OEt

EtO

–EtOOC–COO– O

R

EtO

O

O

O

(16.43)

O O R΄

+e–

R

R΄ –

EtO

–EtOOC–COO– O

R

R

O

Related to this is the reductive elimination of acetate from 1,2-diacetoxy-l,2-diphenylethylene in DMF to diphenylacetylene [157]. In the presence of phenol, diphenylacetylene is further reduced to the stilbene. 3. Sulfonates and Related Esters The reductive conversion of an aliphatic alcohol to an alkane may be carried out in high yield by preparing first the methanesulfonate of the alcohol after which the methanesulfonate is reduced at a lead cathode in DMF containing a tetrabutylammonium salt as supporting electrolyte, Equation 16.44 [158]: R

OSO2Me

+2e–, +H+

(16.44)

RH + MeSO3–

Similarly, tert-butyl benzenesulfenate is reduced to isobutane and the benzenesulfenate anion; the reaction proceeds via a stepwise dissociative mechanism [159]. In contrast, the reduction of benzyl benzenesulfenate involves cleavage of the S–O bond. Cyclic sulfates of unactivated diols such as 2,3-butanediol are reduced in DMF/Et4NC1O4 at a mercury electrode to 2-butene (cis–trans, 5:95) in moderate yields; minor amounts of several unidentified compounds were observed by GLC, Equation 16.45 [160]. H3C

CH3

O

O S

O

~1F –SO4 2–

H3C

CH CH 39–49%

(16.45)

CH3

O

4. β-Arylthio- and β-Alkylthioesters and -Alcohols The reductive elimination of a hydroxyl and a phenylthio group from a β-hydroxysulfide is a key step in a convenient method for the conversion of carbonyl compounds to the corresponding alkenes, Equation 16.46 [161]. If the alcohol is first converted to the methanesulfonate, reduction leads instead to the cyclopropane [162]. R΄

R΄ O R˝

© 2016 by Taylor & Francis Group, LLC

1) PhSCH2Li 2) H3O+

R΄ +2e– –OH–, –PhS–

HO R˝

SPh

CH2 R˝

(16.46)

580

Organic Electrochemistry

A modification of this process in which the carbonyl compound is treated instead with (PhS)2CHLi or PhS(CH3O)CHLi has been used to convert an aldehyde to the next higher homolog, Equation 16.47 [163]. The resulting enol ether or thioenol ether may be transformed into an aldehyde by hydrolysis or reaction with mercuric chloride. R΄ O

1) (PhS)2CHLi

R΄

+2e– –OH–, –PhS–

HO

2) H3O+

R˝

SPh SPh

R˝

SPh

R΄

CH

H3O+

R˝

R΄ CHO

(16.47)

R˝

1,2-Diphenyl-2-(phenylthio)ethyl acetate may be converted to stilbene by reductive elimination of PhS− and AcO − in DMF in a similar way. However, in this case the product, stilbene, is more easily reduced than the substrate and the stilbene radical anion thus formed enters a catalytic cycle in which an electron is transferred to the substrate, Equation 16.48 [164]. Ph

Ph

Ph

Ph – +e–, –PhS–, –AcO– (3 steps)

+e– PhS

OCOCH3 Ph

CH

PhS CH

+e–

Ph

OCOCH3 Ph

CH

Ph –

CH

Ph

Ph – +e–, –PhS–, –AcO–

OCOCH3

PhS

(3 steps) –Ph

(16.48)

Ph

Ph

CH

CH

Ph

PhS

etc.

OCOCH3

5. β-Nitroesters Acylated β-nitroalcohols, prepared by condensation of an aldehyde with a nitroalkane followed by acetylation, may be reduced to alkenes in moderate to good yields, Equation 16.49 [165]. R

R˝

R˝ NO2

R΄O

+2e– –NO2 – , –R΄O–

(16.49)

RCH R΄˝

R΄˝

6. Phosphates Aromatic hydroxyl groups may be removed by conversion first to an aryl diethyl phosphate that is then electrochemically cleaved in DMF to the hydrocarbon and diethyl phosphate anion, Equation 16.50 [166]. O ArO

P OEt

O OEt

+2e–, +H+

ArH + –O

P

OEt

(16.50)

OEt

Simple triaryl phosphates are cleaved in a similar fashion [109,167]. This is in contrast to tris (4-nitrophenyl) phosphate that owing to the electron-withdrawing nitro groups undergoes cleavage

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Cleavages and Deprotections

at the dianion stage accompanied by dimerization [168] similar to what is observed during the reduction of 4-cyanodiphenyl ether [169] (see Section VI.A.10). 2-Haloethyl diphenylphosphates are cleaved reductively to diphenylphosphate [170]. Protection/ deprotection based on these types of compounds has been employed during nucleotide synthesis using the triester method to achieve selective cleavage of one of the ester groups. The 5-trityl protected nucleoside was first treated with 2,2,2-trichloroethylphosphoric acid dichloride and then with 2,2,2-tribromoethanol. By electrochemical reduction in MeCN, the tribromoderivative could be selectively removed, whereas the detritylation could be carried out with 1% trifluoroacetic acid in methylene chloride [171]. 7. Hydroxyketones The hydroxy groups in α-hydroxyketones [172,173], as, for example, in 16-hydroxy-17-ketosteroids, may be reductively cleaved according to Equation 16.51. OH

R

H

R +2e–, +2H+

(16.51)

–H2O R'

O

R'

O

8. Hydroxysulfones Alkenes result also from the reduction of β-hydroxysulfones [174,175] that are easily prepared from the corresponding esters by reaction with ArSO2CH2MgI followed by reduction with NaBH4 Equation 16.52. This route provides a convenient way to convert an ester to an alkene [174]: O RCOOCH3

1) ArSO2CH2MgI 2) H3O+

RCCH2SO2Ar

NaBH4 MeOH

(16.52)

OH RCHCH2SO2Ar

~4F

RCH

CH2

9. Alcohols The C–OH bond in an α,β-unsaturated alcohol, such as allylalcohol and cinnamylalcohol, may be reductively cleaved by direct electrolysis [176,177] or electrocatalytically at a platinized Pt electrode in acidic medium [178] to the alkene and further to the corresponding hydrocarbon, Equation 16.53 [176–178]: Ph

CH

CH

CH2OH

+2e–, +2H+ Ph –H2O

CH

CH

CH3

+2e–, +2H+

Ph

CH2 CH2 CH3

(16.53)

Another approach includes reduction under acidic conditions in the presence of iodide ion; the in situ generated iodo compound is then reduced at a mercury cathode in a reaction that includes a shift of the double bond to the terminal position. Crotyl alcohol may be reduced in this way to a 95/5 mixture of 1- and 2-butene in a total yield close to quantitative [179]. 10. Acyclic Ethers and glycosides The radical anion of diphenyl ether undergoes C–O cleavage in DMF to the phenyl radical and a phenoxide ion with a first-order rate constant equal to 4·105 s−1 [180]. The phenyl radical is

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Organic Electrochemistry

subsequently reduced to the anion and then protonated. The reduction of alkyl aryl ethers follows the same mechanism [181]. This is in contrast to the radical anions of 2- and 4-cyanodiphenyl ether that were found to dimerize. In the case of 4-cyanodiphenyl ether, the resulting dimer dianion undergoes loss of phenoxide giving the 4,4′-dicyanobiphenyl that is further reduced [169]. The coupling pattern of 2-cyanodiphenyl ether radical anion includes three dimer dianions. Allyl aryl ethers such as allyl phenyl ether are reduced in a Ni(bpy)32+ catalyzed reaction to the phenols in high yield [182]. A chlorine atom in phenyl group is not affected. Allyl alkyl ethers are cleaved similarly. Since the allyl aryl ethers are easily obtained from the corresponding phenols by reaction with an allyl halide under basic conditions, the reaction is useful in protection/deprotection chemistry. Similarly, the 4-picolyl derivative of, for instance, l-tyrosine is easy to prepare and to remove electrolytically [183]. Cleavage of an ether bond is observed also during the 6F+6F reduction of 4-alkoxy-1,3-dinitrobenzenes under acidic conditions. The first 6F reduction process is accompanied by hydrolysis of the OR substituent leading to 2-amino-4-nitrophenol that is further reduced to 2,4-diaminophenol in the second 6F process [184]. Benzyl methoxymethyl ethers carrying electron-withdrawing substituents such as p-cyano and o- or p-methoxycarbonyl in the benzyl group are reduced in MeOH to the corresponding toluenes in fair to good yields in reactions that include elimination of methoxide ion at the radical anion stage, Equation 16.54 [185]. In addition, the ester group is reduced to the aldehyde and alcohol stage when the reduction is carried out in the presence of acetic acid. CH2OMe –

CH2OMe +e–

CH3

–MeO–

COOMe

COOMe

(16.54)

+e–, +H+

COOMe O

OMe

Trityl [186] and cinnamyl [187,188] groups are useful in alcohol protection/deprotection chemistry as they are easily removed electrochemically and similar to what was observed for the cleavage of esters [148], cinnamyl ethers are reduced in preference to allyl ethers. Related to these groups is the tritylone group (9,10-dihydro-10-oxo-9-phenyl-9-anthracenyl) and selective protection/deprotection of alcohols may be achieved by using a combination of the tritylone and p-cyanobenzyl groups. Tritylone alcohol (TrOH) reacts preferentially with primary hydroxy groups and is the more easily reduced group. This has resulted in the following reaction sequence for 1,4-pentanediol, Equation 16.55 [189]. By adjusting the reduction potential, either one or both of the protecting groups may be removed. CN

OH OH

OH O

OH

OTr –1.4 V vs Ag/AgCl

O OTr

NC –2.1 V vs Ag/AgCl

NC

CH2Br

(16.55) O OH

OH OH

Glycosidic bonds may be reductively cleaved in DMF by using a redox catalyst such as anthracene, Equation 16.56 [190]. The radical is further reduced by the anthracene radical

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583

Cleavages and Deprotections

anion and then protonated. The redox catalysis may be carried out intramolecularly by attaching acetophenone covalently to the aliphatic chain. O

O HO

HO OCH2(CH2)7CH2CH3 O

O

O

–OCH

2(CH2)7CH3CH3

O

(16.56)

– OH

OH

11. 1,3-Dioxolanes and 1,3-Dioxanes The protection of carbonyl groups by conversion to 1,3-dioxolane derivatives is a classic reaction in carbonyl chemistry. 1,2- and 1,3-Diols substituted with, for instance, a p-nitrophenyl group give 1,3-dioxolanes and 1,3-dioxanes that are easily reductively cleaved. Examples include 4-(4′-nitrophenyl)-1,3-dioxolanes, 5-nitro-1,3-benzodioxane and 7-nitro-1,3-benzodioxane, Equation 16.57, here illustrated by the deprotection of cyclohexanone [191].

O

O O

+e–

O2N

O O2N –

O

(16.57)

O–

O –O

+e–, +H+

N+

O2N

–O

12. Oxiranes Oxiranes carrying phenyl substituents are reduced with ring opening in DMF to a variety of products, including phenyl substituted alcohols, alkenes, and alkanes, the composition of the product mixture being dependent on both the number and position of the phenyl groups [192]. A typical example is given by Equation 16.58. Slightly different product compositions are observed when the reduction is carried out by indirect electrolysis. Ph

Ph

Ph

Ph

Ph 2F

+

O Ph 70%

+ HO

Ph 22%

Ph

(16.58)

8%

Benzoyloxiranes are reduced in MeCN to the corresponding aldoles; addition of acetic acid prevents the retro-aldol reaction catalyzed by the base generated during reduction [193]. The ring opening is analogous to that observed in the reduction of 2-alkoxyacetophenone to acetophenone. Acetyloxiranes react similarly. If the reduction is carried out in aqueous EtOH/Me4NBr in the

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Organic Electrochemistry

presence of catalytic amounts of HBr, the product is the corresponding β-acetoxyketone, which has been used in the reductive conversion of epoxypregnenone (16α,17α-epoxypregn-5-en-3β-ol-20one) to acetoxypregnenone [194].

B. OXIDATIVE C–O CLEAVAGES 1. Enol Esters and Related Compounds The radical cations of enol derivatives such as carboxylic acid esters, carbonates, carbamates, and anhydrides undergo cleavage of the O–CO bond, resulting in the formation of a benzofuran derivative as the final product, here shown for a mixed carbonate, Equation 16.59, Mes = mesityl [195]. The rate constants for the cleavage step were determined by fast sweep CV in MeCN and found to be in the range (2–5) · 103 s−1. O +

O Mes

O OCH2Ph

Mes

–e–

R

O

Mes

OCH2Ph Mes

–PhCH2OCO

R

(16.59) +

Mes

O

+ Mes

R

O

[1,2]-Me

R

–H+

R Mes

Mes

2. Ethers It was observed early that anodic oxidation of hydroquinone mono- and dimethyl ethers in dilute sulfuric acid resulted in the formation of benzoquinone [196]. Similarly, the oxidation of dimethoxydurene in acetonitrile gave duroquinone rather than the side-chain substitution product that might have been expected [197] (see Chapter 23). The reaction sequence shown in Equation 16.60, including the elimination of R+, was offered to explain the observation. OR

OR

O

+ –e– –R+

OR

OR O

OR O

–e– –R+ + OR

© 2016 by Taylor & Francis Group, LLC

(16.60)

O

585

Cleavages and Deprotections

Related to this is the electrochemical oxidation of vitamin E model compounds [198,199]. The reaction has been shown to proceed as illustrated by Equation 16.61 with the key step being the waterassisted cleavage of a C–O bond (see also Chapter 40). OH

O

O –2e–, –H+

H2O

+2e–, +H+

–H+

O

O +

(16.61) O HO

The substituted benzoquinone product may be reduced back to the 2,2,5,7,8-pentamethylchroman6-ol starting material in MeCN/AcOH containing sodium acetate [199]. The electrochemical oxidation of benzyl ethers, ArCH 2OR, leads to the corresponding benzaldehydes, ArCHO, and alcohols, ROH, in a process that includes cleavage of the C–O bond [200]. The reaction has been put to use in the oxidative cleavage of benzyl ether protected alcohols. Anisyl ethers are preferred owing to their relatively low oxidation potential, 1.65 V versus SCE [201]; the yields of the deprotected alcohols were typically close to 90%. Deprotection may preferably be carried out in MeCN containing a base such as 2,6-dimethylpyridine or NaHCO3 using a redox catalyst such as tris(p-bromophenyl)amine; in that case, the operating potential may be lowered by ~550 mV relative to that for the direct electrochemical oxidation [202]. Unsubstituted benzyl ethers cannot be cleaved by the tris(p-bromophenyl)amine radical cation. However, triphenylamines containing bromine atoms not only in the three p-positions, but also in one or more o-positions have oxidation potentials high enough to cleave the simple benzyl ethers [203]. Thus, the oxidation potential of the catalyst may be fine-tuned by adjustment of the number of bromine atoms in the triphenylamine. Alternatively dioxo-bridged binuclear manganese complexes may be used as mediators in MeCN [204]. Also, the p-methoxyphenyl group has been used successfully in alcohol protection/deprotection; in this case, the deprotection step includes the oxidation at platinum electrodes in a MeCN/water (9/1) mixture containing NaHCO3/NaClO4 [205]. Carbonyl groups protected by conversion to 4-phenyl-l,3-dioxolanes may be deprotected by oxidation in MeCN/pyridine using N-hydroxyphthalimide as a redox catalyst [206]. Advantage has been taken of electrochemical deprotection by oxidative C–O cleavage in activating self-assembled monolayers as shown schematically in Scheme 16.7 [207,208]. In this way, longchain aldehydes (left) [208] and carboxylic acids (right) [207] attached to a gold electrode could be exposed for further reaction by a single CV scan.

VII. CARbON–SULFUR bONDS The reader interested in the electrochemical cleavage of the C–S, C–Se, and C–Te bonds is referred also to Chapter 27.

A. REDUCTIVE C–S CLEAVAGES 1. Sulfonium Ions The first step in the electrochemical reduction of a sulfonium ion is a one-electron process accompanied by cleavage of a C–S bond [209–212] similarly to what is observed for other organic “onium”

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586

Organic Electrochemistry OH

O

H

O

O

O

O

+0.2 V

+0.7 V

S

S

S

Au

SCHEME 16.7

O

HO

S

Au

Oxidative deprotection/activation of self-assembled monolayers.

ions, Equation 16.62. The one-electron transfer and the bond breakage may proceed in a stepwise or concerted manner dependent on the structure of the sulfonium ion [211,213] (see Chapter 14 for details): R1 S+

R2

+e–

R1

S

R2 + R

(16.62)

R

The reversible one-electron reduction of sulfonium ions to the corresponding free radicals has been observed only for diarylalkylsulfonium ions carrying one or more nitro groups in one of the aryl groups and only at low temperature (−40°C) [214]. Amalgams may be formed by reduction at mercury electrodes [53,215]. For sulfonium ions in which R, R1, and R2 in Equation 16.62 are not all identical, the question of the relative ease of cleavage of the three bonds has received considerable attention [211,212,216–218]. For a series of monophenylalkylsulfonium ions, reduction resulted in cleavage of the weaker S-alkyl bonds in the order benzyl > secondary > primary > methyl > phenyl [209,216] reflecting the stability of the resulting radical. In contrast, the stronger S-aryl bond is preferentially cleaved during reduction of diphenylmethylsulfonium ions [212,216,218]. Results from indirect electrolysis have shown that the ratio of S-aryl/S-alkyl cleavage depends on the reduction potential of the reducing agent, and it was concluded that the differences between the reductive cleavage of mono- and diarylsulfonium ions are direct consequences of the structures of the sulfuranyl radicals and the bond dissociation energies of S-alkyl and S-aryl bonds [217]. The radical, R•, formed during the cleavage reaction has been detected by spin trap experiments [214,219,220] and dimers derived from R· have occasionally been detected [54,221]. Radicals generated by electrochemical reduction of sulfonium ions have been utilized in surface derivatization of glassy carbon electrodes [222] (see also Chapter 42). During coulometry or preparative electrolysis, reduction may proceed as a 2F process including the further reduction of R• to R− that finally may be protonated to RH. However, RH may arise also via a hydrogen-abstraction reaction between R• and a solvent component. When the reduction is being studied by redox catalysis, products resulting from coupling between R• and the catalyst radical anion are observed [223,224].

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587

Cleavages and Deprotections H3C S+

CH2Ph

+e–

H3C

S

CH2 + PhCH2

H3C +e–

PhCH 2–

PhCH2 H3C

H3C S+

H3C

H3C

SCHEME 16.8

– CHPh + PhCH3

S+

CH2Ph + PhCH2–

Ylid formation resulting from reductive cleavage of a benzylsulfonium ion.

The product distribution resulting from reduction of triphenylsulfonium ion depends on both the electrode material and the potential [212,215]. Reduction of triphenylsulfonium bromide at a mercury cathode at the potential of the first wave gives quantitatively diphenylsulfide and diphenylmercury. At potentials lower than the second wave, diphenylsulfide and benzene are formed, provided low-substrate concentrations are used. With increasing substrate concentrations, the yield of diphenylmercury increases at the expense of benzene. Reduction in aqueous solution at an aluminum cathode produces diphenylsulfide and minor amounts of benzene; addition of DMF increases the yield of benzene indicating that benzene is formed (at least partly) in a hydrogen-abstraction reaction [212]. The base produced during reduction may react with sulfonium ions carrying a hydrogen atom in the α-position resulting in formation of the corresponding ylid (Scheme 16.8) [54,225,226]; the ylid could be trapped by, for instance, benzaldehyde to give an epoxide by the Corey–Chaykovsky reaction. Ylid formation has been attributed to be the origin of the inhibition of the reduction process observed at graphite cathodes [116]. Attempts to utilize the anion, R−, in a Michael condensation with acrylonitrile resulted in only moderate yields of the expected products [227]. 2. Thiol Esters The product distribution resulting from reduction of thiol esters in DMF/LiClO4 [228], DMF/ Bu4NI [229], or MeCN/Et4NBF4 [138] is dependent on the nature of the two groups attached to the –CO–S– function [138,228] and also on the supporting electrolyte cation [228,229]. The major products observed in DMF/LiClO4 [228] may be summarized as follows: (a) R–CO–S–R gives R–CO–SH, (b) R–CO–S–Ar gives R–CO–NMe2 (with the NMe2 part originating from DMF), (c) Ar–CO–S–R gives Ar–CO–CO–Ar, and (d) Ar–CO–S–Ar gives Ar–CO–CO–Ar, where R in all cases is an alkyl-type group and Ar an aryl group. In DMF/Bu4NI [229] reduction of Ph–CO–S–Ar and Ph–CO–S–R proceeds all the way to the 1,2-diphenylacetylene. If oxygen is present during the reduction, R–CO–OH is formed in a reaction with super oxide ion [140]. Thus, it appears that the primarily formed radical anion may undergo cleavage of either the CO–S bond or the S–Y bond (Scheme 16.9), where Z and Y are the alkyl/aryl groups. When Z = Ar, O + Path a

O

Y Z

S

Y

O

Path b

+ Y Z

SCHEME 16.9

–S

Z

S–

Competition between CO-S and S-Y cleavage during reduction of thiol esters.

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588

Organic Electrochemistry

path  (a)  is  preferentially followed and dimerization of the radical then leads to the diketones, whereas Z = R results in products preferentially derived from path (b). However, another pathway leading also to α-diketones includes the dimerization of the radical anions followed by elimination of YS−. This was found to the likely mechanism for reduction of benzenecarbodithioic esters to diphenylacetylenes [229]. Cleavage of a C–S bond in the initially formed radical anion has been shown to be the first chemical step following electron transfer in the electrochemically induced rearrangement of S,S-diaryl benzene-l,2-bis(carbothioates) to 3,3-bis(arylthio)isobenzofuran-1(3H)-ones in DMF by Equation 16.63 [230]. Complete conversion required less than 0.1F indicating that an SRN1-type mechanism with elimination of ArS− at the radical anion stage is followed. O O SAr ~0.1F

(16.63)

O SAr SAr ArS

O

The S-carbobenzoxy group has been used as a protection group for cysteine and could be removed by reductive cleavage in 72% yield [231]. The differences in reduction potentials for N-carbobenzoxybutylamine, O-carbobenzoxybutanol, and S-carbobenzoxycysteine make possible the selective removal of the protective groups. Dialkyl, alkyl aryl, and diaryl trithiocarbonates are reduced in DMF to the corresponding radical anions followed by loss of thiolate ion [232–234]. The rate constant for the cleavage was found to vary linearly with Hammett’s substituent constant σ; the radical anions of the mono- and dithiocarbonates having a >C=O group were cleaved faster than those having a >C=S group. The effect of structure on the electron transfer kinetics was rationalized by using a multisphere solvation model for the solvent reorganization within the Marcus–Hush theory. When the reduction was carried out in the presence of a suitable alkylating agent such as dimethyl sulfate or ethyl iodide, the corresponding 1,1,2,2-tetrakis(alkylthio)ethene was obtained, Equation 16.64 [232,234]. S

RS

1) +4e–

2 PhS

SR + 2PhSR + 4X–

2) +4RX

SR

RS

(16.64)

SR

3. Thioethers (Sulfides) The C–S bond in dialkylsulfides is difficult to cleave by electrochemical reduction; however, cleavage at a silver cathode has been reported [235]. In contrast, most alkyl aryl and diaryl sulfides are reduced in DMF [236] at a low potential (~ –2.5 V vs. SCE), or in THF at low temperature [237], to ArS− and R− (or RH) in an altogether 2F process via the one-electron reductive cleavage of the C–S bond, Equation 16.65: Ar

S

R

+e–

Ar

S– + R

(16.65)

For Ar = 2-naphthyl, it was observed that the mode of cleavage in THF could be controlled by the temperature [238]. At 20°C, the reaction proceeded predominantly via cleavage of the alkyl-S bond as in Equation 16.65, whereas cleavage of the aryl-S bond at the dianion stage was observed at −78°C. When naphthalene was used as a redox catalyst in THF at low temperature, the cleavage

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Cleavages and Deprotections

of Ar–S–R could be carried out at a potential approximately 500 mV higher than that for the direct process. Under those conditions, not only RH was isolated, but also PhSSPh [239]. The bond-breaking details have been investigated by redox catalysis for phenyl triphenylmethyl sulfide [240–242] and a series of p-substituted phenyl triphenylmethyl sulfides [241,242]. For the cleavage of diphenylmethyl p-methoxyphenyl sulfide, the mechanism includes the protonation of the diphenylmethyl anion by the starting material (self-protonation) [243]. The data obtained for strongly electron-withdrawing substituents were in agreement with a stepwise mechanism; a gradual transition toward a pathway including a loose radical anion was observed in passing through the unsubstituted compound to the p-methoxy-substituted (see Chapter 14 for details). Phenacyl phenyl sulfide is reduced in an aqueous-alcoholic solvent mixture to acetophenone and thiophenol [173]. This particular reaction has been put to use in the cleavage (deprotection) of 2-acyl-3-aminothiophene derivatives in DMF containing phenol as the proton donor, Equation 16.66 [244]. The starting material was easily prepared by reaction between an α-haloketone and sodium cyanophenyldithioacetate. H2N

Ph

H2N

Ph

– + Ar +2e , +2H

Ar S

–ArCOCH3

S

(16.66)

Ar S

O

SH

O

O

Deprotection of thiols in general may be accomplished by reduction of the corresponding benzyl thioethers, the prominent example being the deprotection of the thiol group in cysteine by reduction at a platinum cathode in liquid ammonia [245] or a mercury cathode in MeOH [246]. Other thiolprotecting groups that may easily be removed by electrochemical reduction include trityl [231] and 4-picolyl [183,247]. The reduction of gem-disulfides such as 3,3-bis(methylthio)-1-phenylprop-2-en-1-one leads to cleavage of the C–S bond and elimination of CH3S−, Equation 16.67 [248]. The free radical was reported to polymerize. In the presence of carbon dioxide, substitution of the CH3S group by COO − was observed. O

O SCH3

Ph H

SCH3

+e–

O SCH3

Ph

SCH3

Ph

+ C H3S– H

SCH3

(16.67)

H

Related to this is the reductive cleavage of arylthiophenylacetylenes in DMF in the presence of phenol; the C(sp)–S bond is cleaved resulting in the formation of the ArS− anion and the Ph–C≡C· radical [249]. The reaction was also studied by redox catalysis. It was concluded that the bond cleavage was the rate determining step. An exotic example of a reductive C–S cleavage is the removal of 2,2′-oxydiethanethiol and similar fragments from o-carboranyl derivatives such as that shown in Scheme 16.10 [129] (the noncarbon atoms in the icosahedron are all boron). However, the fate of the 2,2′-oxydiethanethiol fragment is not clear, but it was observed that in the presence of the larger 2,2′-(ethane-1,2-diylbis(oxy)) diethanethio ring system, the reductive elimination of the ring system was accompanied by loss of a boron fragment from the o-carborane structure. 4. Sulfoxides The electrochemical reduction of DMSO at a platinum cathode and in the presence of Bu4NBF4 leads to C–S and S–O cleavage and the formation of MeS− and Me2S, respectively [250].

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Organic Electrochemistry

S C O S

SCHEME 16.10

C

Example of an o-carboranyl derivative that undergoes reductive C-S cleavage. Ar1 Ar1 O

S

+e–

O

Ar2

S

S–

O

Ar2

S

S

O–

S

O–

Ar2 Ar1

Ar1 S + –O Ar2

Ar2

Ar1

O

Ar2

Ar1

Ar1

Ar1

Ar1

Ar2

+e–

S

O–

Ar2

Ar1SO2– + Ar2H O–

or Ar2SO2–

Ar2

SCHEME 16.11

O–

S

Ar2

–O

S

Ar1

+ Ar1H

Reductive disproportionation of diarylsulfoxides.

Diarylsulfoxides undergo disproportionation to the corresponding sulfide and sulfone upon reduction in aprotic solvents followed in most cases by reductive cleavage of the sulfone (see Section VII.A.5) [251]. The disproportionation reaction was suggested to proceed as shown in Scheme 16.11. Reductive cleavage of β-ketosulfoxides is a practical way of preparing methylketones [252,253]; ω-(methylsulfinyl)-p-methoxyacetophenone, for example, gives acetophenone (74%) upon reduction at a mercury cathode, Equation 16.68 [252]. O O

H3CO H2C

+2e–, +H+

S

O H3CO

+ CH3SO– CH3

(16.68)

CH3

The major products resulting from the reduction of vinyl aryl sulfoxides in DMF/Bu4NBF4 are the corresponding arenesulfinate anions resulting from disproportionation (see Scheme 16.11) and subsequent cleavage of the ArSO2–CHCHR bond [254]. Direct cleavage of the ArSO–CHCHR bond has been reported to take place for vinyl phenyl sulfones that carry an α-CF3 substituent. The resulting phenylsulfanolate anion undergoes reduction to the thiophenolate that reacts with the starting material in an altogether very complex reaction scheme [255]. Reduction of (E)-l-methylsulfinyl-l-methylthio-2-phenylethene with excess of phenol includes the selective cleavage of one C–S bond and the formation of (E)-l-methylthio-2-phenyl-ethene, Equation 16.69. The dependence of the CV curves on the phenol-to-substrate ratio and the stereochemistry of the

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Cleavages and Deprotections

reaction supported the mechanism shown including as the second step the elimination of the methylsulfanolate anion [256]. Other α,β-unsaturated sulfoxides react similarly [255,257]. O

O H

Ph

S

S

S

H

CH3

+2e–, +H+

CH3

CH3

H

SCH3

– Ph

S

CH3

(16.69)

–CH3SO– Ph

H

5. Sulfones The electrochemical reduction of sulfones has been the subject of a number of reviews to which the reader is referred for details [209,258]. Simple dialkyl sulfones and bis(alkylsulfonyl)methanes are difficult to reduce electrochemically. An exception is the disulfone derived from 1,3-dithietane that is reduced in a 2F process with ring opening to methylsulfonylmethanesulfinate, Equation 16.70 [259]. +2e–, +H+ O2S

SO2

CH3SO2CH2SO2–

(16.70)

The general reaction scheme for the reduction of alkyl aryl sulfones includes cleavage of the ArSO2–R bond at the radical anion stage; the resulting alkyl radical is further reduced to the anion that is finally protonated to RH (Scheme 16.12) [259–262]. However, when the reduction is carried out at low temperature in THF, the radical anion has a lifetime that is sufficient to allow for the observation of its oxidation back to the substrate during CV [263]. The presence of electron-withdrawing substituents such as p-NO2 and p-CN causes the radical anions to be long-lived and they can easily be detected by CV at room temperature [264]. For sulfones of this type [264,265], cleavage of the Ar–SO2R bond with elimination of the alkane sulfinate anion is the predominant reaction pathway; the cleavage reactions in those cases may be so slow that other follow-up reactions such as protonation and dimerization may take place [266]. Similarly, 2- and 4-(alkylsulfonyl)pyridines and related sulfones undergo Ar–SO2R cleavage [267]. This offers a convenient route to alkylsulfinic acids and thus to alkylsulfones with two different alkyl groups. For long-chain alkyl groups, the basic conditions created during reduction in, for instance, DMF may cause the formation of alkenes by β-elimination [261]. The reduction of seven-membered ring sulfones such as dibenz[b,e]-thiepin-11-one-5,5-dioxide (Scheme 16.13) and the related 11-thione is accompanied by cleavage of the CH2–SO2 bond as expected [268]. An unexpected cleavage reaction was observed for the reduction of 4-phenylthio-tertbutylsulfonylbenzene; in that case, the elimination involved the 2-methylpropane-2-sulfinate anion that subsequently attacked the starting material in a substitution reaction, resulting in the formation of the symmetrical 1,4-bis(tert-butylsulfonyl)benzene [269].

ArSO2R

+e–

ArSO2R –

+e– R–

R

+H+ RH

+ArSO2R – –ArSO2R

SCHEME 16.12

Reductive cleavage of alkyl aryl sulfones.

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ArSO2– + R

592

Organic Electrochemistry O O S

O

SCHEME 16.13

Structure of dibenz[b,e]-thiepin-11-one-5,5-dioxide.

The reduction of diaryl sulfones follows a mechanism similar to that in Scheme 16.12 [260,270–272]. For mono-substituted diphenyl sulfones, the selectivity of the bond cleavage was dependent on the electronic nature of the substituent. For reduction at mercury in MeOH containing Me4NCl/Me4NOH (2/1), it was found that electron-donating substituents resulted in cleavage of the S–Ph bond, whereas the bond between sulfur and the substituted phenyl group was cleavaged exclusively when the substituent was electron withdrawing [271]. Substituents in the o-position facilitated cleavage of the neighboring C–S bond, probably by steric hindrance of the conjugation between the aromatic ring and the sulfone group [271,273]. A special case is 1-(4-biphenylyl)-2-phenylsulfonyl-3,3,3-trifluoropropene that in a +2e−, +H+ process including the cleavage of a C–S bond is converted to Ar–CH=CH–CF3 and PhSO2− [255]. The central hydrogen atoms in bis(arylsulfonyl)methanes are acidic (pKa = 12.5 for bis (phenylsulfonyl)methane) owing to the two neighboring SO2-groups and, accordingly, the mechanism for reduction depends on the availability of protons in the solvent-supporting electrolyte mixture. In a nonbuffered DMF/water mixture (20/80) with Et4NBr as the electrolyte, reduction proceeds according to Equation 16.71 with cleavage of the C–S bond taking place at the dianion stage [259]: ArSO2CH2SO2Ar

+e–, +e–

ArSO2CH2SO2Ar2–

H2O –OH–

ArSO2CH3 + ArSO2–

(16.71)

This is in contrast to the mechanism observed in strictly aprotic solvent in which a self-protonation scheme is observed (Scheme 16.14). In both cases, the alkyl aryl sulfone produced may be further reduced, but the benzenesulfinate anion is not electroactive in the potential range applied. Self-protonation reactions are observed also for the reductive cleavage of allylic and benzylic sulfones in aprotic solvents [274], whereas reduction of vinyl sulfones results in cleavage to the corresponding alkenes and benzenesulfinate [114,255,275]. Analogous to this is the reductive cleavage of the C–S bond in β,β-bis(methylsulfonyl)styrene to β-(methylsulfonyl)styrene [276]. In mixed disulfones that have both the ArSO2CH2R and the ArSO2CH2CH=CH2R structural features, the allylic benzenesulfonyl group is preferentially cleaved; the selectivity is higher when indirect electrolysis is being used [277]. For cyclic 1-cycloalken-1-yl phenyl sulfones, the cleavage reaction is accompanied by isomerization into the corresponding allyl sulfones triggered by the electrogenerated base [278]. 1-Methylthio-1-p-tolylsulphonyl-2-arylethenes are reduced in a fashion analogous ArSO2CH2SO2Ar

+e–

ArSO2CH2SO2Ar –

ArSO2CH2SO2Ar – + ArSO2CH2SO2Ar [ArSO2CH2SO2Ar]H + ArSO2CHSO2Ar– [ArSO2CH2SO2Ar]H

SCHEME 16.14

+e–,+H+

ArSO2CH3 + ArSO2–

Reductive cleavage and self-protonation of a bis(arylsulfonyl)methane.

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Cleavages and Deprotections

to the reduction of l-methylsulfinyl-l-methylthio-2-phenylethene described earlier, Equation 16.69, in this case with the elimination of the p-methylbenzenesulfinate anion as the second step [279]. The ability to act as a leaving group during the reductive cleavage of diaryl sulfones or alkyl aryl sulfones is benzyl ~ allyl > alkyl > phenyl [260,261,271,272]. The reductive cleavage of benzene substituted with more than one alkyl- or arylsulfonyl group has been studied in detail [280–284]. The products resulting from reduction of the disubstituted compounds depend on the substitution pattern and on whether the substituent is an alkyl- or arylsulfonyl group. Reduction of 1,2-bis(alkylsulfonyl)benzenes [282,285] and 1,4-bis(alkylsulfonyl)benzenes [283] leads via the formation of the radical anions to the formation of self-alkylation products in addition to the expected cleavage products, here shown for a 1,2-bis(alkylsulfonyl)benzene, Equation 16.72. SO2

R

+2e–

R

H+ (solvent)

SO2

SO2

SO2

R

R

+

(16.72)

SO 2 –

–RSO2– R

The reduction of 1,2-bis(phenylsulfonyl)benzenes proceeds in a slightly different manner. The free radical resulting from elimination of the benzenesulfinate anion undergoes hydrogen abstraction to the diphenyl sulfone and cyclization followed by elimination of a hydrogen atom giving the dibenzothiophene dioxide (Scheme 16.15) [280]. The reduction of 1,3-bis(alkylsulfonyl)benzenes and 1,3,5-tris(alkylsulfonyl)benzenes proceeds completely different; the radical anions were found to dimerize to a dimer dianion [282,284] similarly to what is observed for other benzenes substituted in the 1- and 3-positions with strongly electron-withdrawing groups (see Chapter 17). The radical anions of tetra- and pentasulfones are stable at the time scale of slow sweep CV, whereas those of the hexasulfones undergo reductive coupling accompanied by the loss of sulfinate anions, Equation 16.73, in addition to the cleavage reaction [281,282]. SO2R

SO2R

RO2S +2e– –2RSO2–

SO2R

RO2S

SO2R

SO2R

RO2S

RO2S

SO2R

SO2 Ph

+e–

SO2

Ph SO2

SO2

SO2R

Ph

Ph SO2

Ph

Ph Cyclization and hydrogen elimination SO2

Reductive cleavage of 1,2-bis(phenylsulfonyl)benzene.

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SO2R

–PhSO2–

SO2 SO2

RO2S

Ph –

Hydrogen abstraction

SCHEME 16.15

RO2S

SO2R

RO2S

(16.73)

594

Organic Electrochemistry

The reductive cleavage of β-ketosulfones [286], RCOCHR′SO2R″, in DMF at a mercury cathode is a practical way of preparing alkyl ketones and proceeds analogously to the reduction of β-ketosulfoxides, Equation 16.68. The anode reaction is the oxidation of R″SO2− to R″SO3− and an undivided cell may then be used. The sulfamoyl group in benzenesulfonamides carrying strongly electron-withdrawing substituents may be eliminated as sulfur dioxide and ammonia in a pH independent step by reduction at a mercury cathode in aqueous solution [287]; thus, cleavage of the C–S bond occurs before cleavage of the S–N bond. The sulfamoyl groups in benzene-1,3,5-trisulfonamide and benzene-1,3-disulfonamide may be eliminated one-by-one in this way. 6. Thiocyanates The electrochemical reduction of p-CN and p-NO2 substituted benzyl thiocyanates in MeCN leads exclusively to cleavage of the ArCH2–SCN bond (Scheme 16.16, path (a)) and the formation of the corresponding bibenzyls [288]. Reduction in the presence of an excess of phenol gave the toluene, ArCH3, which was taken as an indication that the bibenzyl was formed in a substitution reaction between ArCH2−, formed by reduction of the benzyl radical, and the starting material and rather than by dimerization of the benzyl radicals. In contrast, benzyl thiocyanates substituted in the p-position with MeO, Me, H, Cl, and F gave the corresponding mono- and disulfides together with the toluenes resulting from cleavages by both path (a) and (b) in Scheme 16.16. The analysis of the voltammetry data showed that a stepwise mechanism including a distinct radical anion was followed for p-NO2 substituted compound, whereas the compounds substituted with the weaker electron-withdrawing p-CN and all the electron-donating substituents were following a concerted cleavage mechanism. Theoretical DFT data (B3LYP 6–31G+(d,p)) showed the existence of stronger interactions between the produced fragments for path (b) than for path (a); this seems to account for regioselectivity of the cleavage. The effect of a strongly electron-withdrawing group is observed also for the reduction of other thiocyanates. For instance, for 1-phenyl-2-thiocyanatoethanone in which the thiocyanato group has the electron-withdrawing ArCO in the α-position, reduction proceeds with cleavage of the ArCOCH2–SCN bond, Equation 16.74 [173] analogously to the cleavage shown in Equation 16.66. In most other thiocyanates, reduction results in loss of a cyanide ion [289] and macroscale reductions at a mercury cathode gives the mercaptans. O

O SCN

Ar

+2e–, +H+

(16.74)

+ SCN–

Ar

CH2

CH3

B. OXIDATIVE C–S CLEAVAGES 1. Thiol Esters S-tert-butyl thioates can be deprotected by indirect electrochemical oxidation using the bromide/bromine redox system as mediator, Equation 16.75 [290]. By direct anodic oxidation, 4-methoxyphenylthiomethyl

Path a X

CH2 S

CH2 + S CN–

X

CH2 S + CN–

CN – Path b

SCHEME 16.16

X

Competitive cleavages observed during reduction of benzylthiocyanates.

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Cleavages and Deprotections

esters are readily converted to carboxylic acids, Equation 16.75 [290]. The initially formed radical cation cleaves on reaction with water to the acylated hemiacetal and further to the acid. –2e– Br+

O

Br–

Anodic oxidation 10–11F

O R

R

O

(16.75)

R

MeCN/H2O

MeCN/H2O OH

S

OCH2SAr

Triarylamine radical cations may also be used in an indirect electrochemical oxidative cleavage of thiol esters [291]. 2. Thioethers (Sulfides) Anodic oxidation of thioethers usually results in the formation of sulfoxides, sulfones, or sulfonates (see Chapter 27). However, if R in ArSR is a good leaving group such as benzyl or trityl, oxidation in MeCN may lead to cleavage of the C–S bond. Phenyl trityl sulfide, for example, is oxidized to diphenyl disulfide (48%) and triphenylmethanol (52%) (Scheme 16.17) [292]. [4-Hydroxy-3-coumarinyl]-phenylthiomethanes are oxidized in a similar fashion. However, in this case, the intermediate carbocation is trapped by MeCN in a Ritter reaction, resulting in the formation of N-[4-hydroxy-3-coumarinyl]methylacetamide [293]. Related to these is the lowtemperature oxidative cleavage of thioglycosides such as thioglucosides, thiogalactosides, and thiomannosides that when the reaction is carried out in dichloromethane with tetrabutylammonium triflate as supporting electrolyte leads to the corresponding glycosyl triflate pools [294]. The methylthiomethyl group may be used as a protecting group for alcohols; deprotection may be carried out by oxidation in an undivided cell at a platinum anode in AcOH-containing NaOAc. The resulting acetoxymethyl ether may then be hydrolyzed in weakly alkaline solution back to the alcohol, Equation 16.76 [295]: –e–, –H+ R

SCH3

O

+OH– R

AcOH/AcONa

O

OAc

–AcO–

R

OH

(16.76)

Electrooxidative desulfenylation of Michael-type thiol adducts of activated olefins may be carried out indirectly in an undivided cell at platinum electrodes with EtOH and Bu4NBr as solvent and supporting electrolyte, Equation 16.77, Y = COOEt, COMe, and CN [296]. The alkylthio (or phenylthio) residues are oxidized mostly to the sulfinates. The resulting activated olefins are formed in good to excellent yields. SEt R

CH

CH2 Y

Anodic oxidation 3–4F Br–/EtOH

–e– Ph

S

2 Ph

S

+CPh 3

SCHEME 16.17

Ph

CPh3 Ph +H2O –H+

Ph3C

S

S S

CPh3 + Ph

OH

Oxidative cleavage of phenyl trityl sulfide.

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R

CH

CH

Ph

(16.77)

Y

S +

+CPh

3

596

Organic Electrochemistry

S

+

–2e– –1/4 S8

S

+RCN, +H2O

+

RCONH

–H+

SCHEME 16.18

2

Oxidative cleavage of di-tert-butyl disulfide.

The C–S bonds in gem-disulfides may be cleaved by anodic oxidation in aqueous MeCN to the corresponding carbonyl compounds [297–300]. This has been put to use for the removal of the 1,3-dithian protecting group, Equation 16.78 [298]. The oxidation may be carried out indirectly in MeCN/LiClO 4 containing NaHCO3 by using tris(p-tolyl)amine as the electron transfer agent; carried out this way the yields of the recovered carbonyl compound were almost quantitative [301].

R1

S

R2

–2e–

R2

S

H2O

R1

H2O –H+

R1

OH

R2

S+

S+

S

(16.78)

S (–S[CH2]3S–)n

O+

+

–H

S+

R1

R2

S

3. Disulfides The electrochemical oxidation of disulfides results in most cases in cleavage of the S–S bond. An exception is the oxidation of di-tert-butyl disulfide that undergoes cleavage of the C–S bond according to Scheme 16.18 [302]. The yields of the resulting amides are close to quantitative in most cases.

VIII. CARbON–SELENIUM AND CARbON–TELLURIUM bONDS A.

REDUCTIVE C–Se AND C–Te CLEAVAGES

The electrochemical reduction of diphenyl selenide in MeCN leads to cleavage of the C–Se bond with the formation of PhSe – and Ph– anions [303]; the latter are protonated by MeCN resulting in the formation of CH2CN–. The PhSe – anion may be oxidized to diphenyl diselenide by air and reaction of the diselenide with CH2CN– finally leads to 1-phenylseleno-1-cyano-2-aminopropene in moderate yield in a many step process. Unsymmetrical phenylselenobenzonitriles [304] suffer preferentially cleavage of the Ph–S bond, resulting in the formation of the cyanobenzeneselenolate anion in ~85% yield. Reduction of p-ArC6H4SeC≡CAr [305] in MeCN or DMF at a mercury cathode results in cleavage of the C(sp)–Se bond. Similarly to the corresponding selenium compounds, unsymmetrical phenyltellurobenzonitriles suffer cleavage of the Ph–Te bond, resulting in the formation of the cyanobenzenetellurolate anion [306]. Reduction of p-ArC6H4TeC≡CAr [305,307] in MeCN or DMF at a mercury cathode results in cleavage of the C(sp)-Te bond.

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Cleavages and Deprotections

B. OXIDATIVE C–Se AND C–Te CLEAVAGES Electrochemical oxidation of organoselenium and organotellurium compounds usually does not include C–Se/C–Te cleavages. However, an oxidative cleavage reaction has been reported for cyclic alkyl phenyl selenides such as selenochromanone [308] and for compounds of the general structure R–Te–C≡C–Te–R [309].

IX. CARbON–HALOgEN bONDS A.

REDUCTIVE C–Hal CLEAVAGES

The reductive cleavage of the C–Hal bond is probably the most studied electrochemical cleavage reaction. For that reason, Chapters 24 and 25 have been devoted to various aspects of this important process. In addition, a detailed description of the C–Hal reduction process is included in Chapter 14. In this chapter, we shall restrict ourselves to mention briefly a few cases in which halogens have been used as protecting groups. Vicinal dibromides can be converted into olefins by electrochemical reduction and the reduction potential is shifted to higher values with increasing alkyl substitution. This has been used as selective protection of the less alkylated double bond in dienes. The diene is converted into the tetrabromide with bromine and subsequently deprotected by controlled potential electrolysis (Scheme 16.19) [310]. β-Halogenethyl groups have been used as protecting groups for several types of functional groups. Deprotection can be carried out electrochemically by direct [74,75] and indirect [311] electrolysis. Using a vitamin B12-derivative as mediator, penicillin V β-bromoethyl esters can be deprotected to penicillin V under conditions so mild that even the β-lactam ring is kept intact [311].

X. NITROgEN–NITROgEN bONDS In addition to the reactions discussed later, processes that include cleavage of the N–N bond are found also in Chapters 29 through 31.

OH

(C5H11)2BH

Br Br2 Br

Br

+2e–, –2Br–

Br Br Br

B2H6 Br +2e–, –2Br–

HO Br

HO

SCHEME 16.19

Protection-deprotection sequence for 4-vinylcyclohex-1-ene.

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Organic Electrochemistry

A. REDUCTIVE N–N CLEAVAGES 1. Carboxylic Acid Hydrazides and Azides Activated hydrazides [312], such as isonicotinic hydrazide (isoniazid), are reduced to the amides under acidic conditions in a reaction that includes the cleavage of the N–N bond and the formation of ammonium ion, Equation 16.79. NHNH3+

O

NH2

O +2e–, +2H+

+ NH4+

(16.79) +

N

+

N

H

H

The reduction of isonicotinoyl azide also leads to isonicotinic amide, in this case by elimination of a molecule of dinitrogen; most other azides react similarly [313]. The reaction has been put to use in the synthesis of saturated and unsaturated amino acids by cathodic reduction of azidocinnamic esters [314,315]. Reduction of azides in DMF in the presence of acetic anhydride gives diacetylated amines [316]. In contrast, the reduction of azides carrying a neighboring carbonyl or thiosemicarbazone function [85] leads to cleavage of the C–N bond and elimination of azide ions (see subsection III.A.5 on C–N cleavage). 2. Azines and Related Compounds Organic compounds with the common structural feature >C=N–N< may undergo reductive cleavage of the N–N bond under acidic conditions in reactions that include also saturation of the >C=N– double bond. Examples of this behavior include the reduction of benzalazine [316], phenylhydrazones, [316,317], Equation 16.80, semicarbazones [316], and arylidene benzohydrazides [317,318]: R N

+4e–, +6H+

NHPh

R NH3+ + PhNH 3+

(16.80)

R΄

R΄

In aprotic media, the N–N bond appears to be cleaved at the radical anion stage and the rate of cleavage for a series of N,N-dimethylhydrazones of arylaldehydes was found to be a linear function of the potential for the reversible reduction of the hydrazone. Under the basic conditions generated by the reduction, elimination of the amine part of the hydrazone may take place, resulting in formation of a nitrile [319]. 3. Diazo Compounds The reduction of diazo compounds such as diazoacetophenones [320], 3-diazocamphor, Equation 16.81 [321], and 3-diazoindole [322] in aqueous buffers at pH 6–9 proceeds to the corresponding amines in a 6F process including cleavage of the N–N bond. At pH > 9, reduction takes place in two steps, first to the hydrazone in a 2F process and then to the amine in a 4F process. In other cases, reduction of organic diazo compounds is accompanied by cleavage of the C–N bond (see subsection III.A.6 on C–N cleavage). O +6e–, +6H+

O + NH3

N+

NH2 N–

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(16.81)

599

Cleavages and Deprotections

B. OXIDATIVE N–N CLEAVAGES Electrochemical oxidation of benzaldehyde hydrazones may result in cleavage of the nitrogen– nitrogen bond and formation of the corresponding benzonitrile [323]. Oxidative cleavage of an N–N bond has been observed also during the oxidation of N-nitrosoarylamines in MeCN [324]. In one case, N,N′-dimethyl-N,N′-dinitroso-p-phenylendiamine, the results indicated that cleavage took place at the dication stage by elimination of two nitrosonium ions. In two other cases, the ammonium salt of N-nitrosophenylhydroxylamine (known as Cupferron) and N-nitrosodiphenylamine self-inhibition of the electrode process owing to adsorption was observed.

XI.

NITROgEN–PHOSPHORUS bONDS

A.

REDUCTIVE N–P CLEAVAGES

Dimethylaminotriphenylphosphonium ion, Equation 16.82, and related phosphonium ions undergo cleavage of the N–P bond during electrochemical reduction in MeOH [115]: Ph3P+

NMe2

+2e–, +H+

Ph3P + Me2NH

(16.82)

The outcome of the cleavages observed for bifunctional phosphonium ions is less easy to predict. When both a diethylamino and an ethylthio group are present, the N–P bond is preferably cleaved (in MeCN) and a Me2N–P bond is cleaved in preference to a MePhN–P bond (in water), Equations 16.83 and 16.84 [115]: Ph

SEt P+

Me

Ph +2e–, +H+

NEt2

P

SEt + Et2NH

(16.83)

Me

Ph Ph

N P+

Me

Me +2e–, +H+

Ph

Ph P

NMe2

Me

+ Me2NH

N

(16.84)

Me

The N–P bond in triphenylphosphinephenylimine may be cleavaged electrochemically by both reduction and oxidation [325]. The oxidative cleavage leading to triphenylphosphine oxide and polyanilines is believed to occur at the radical cation stage.

XII. A.

NITROgEN–OXygEN bONDS REDUCTIVE N–O CLEAVAGES

Nitrile-N-oxides are easily reduced to the corresponding nitriles; 2,4,6-trimethylbenzonitrile-N-oxide, for example, may be reduced at –0.7 V (vs. SCE) at pH 6 in aqueous ethanol to the nitrile [326]. The reductive cleavage of the N–O bond in oximes generally occurs prior to the saturation of the double bond, Equation 16.85 [316]. However, in alkaline solution at low temperature, reduction of aldoximes to hydroxylamines may take place to some extent. Similarly, cleavage of an N–O bond has been observed during the reduction of amide oximes [316] and hydroxamic acids [313]. The suggestion that cleavage of the N–O bond takes place for oximes bearing a positive

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Organic Electrochemistry

charge on nitrogen was supported by the observation that, after quaternization, isoxazoles, or 4,5-dihydroisoxazoles may undergo N–O cleavage in neutral solution without saturation of the C=N bond [327]: R

H+

OH

R N+

NOH R΄

XIII. A.

+2e–, +2H+

N+

–H2O

H

R΄

H

R

R

+2e–, +2H+

NH3+

H

R΄

(16.85)

R΄

NITROgEN–SULFUR bONDS REDUCTIVE N–S CLEAVAGES

1. Sulfonamides The widespread interest in the electrochemical cleavage of the N–S bond in sulfonamides has been driven primarily by the applications of the process in the protection–deprotection chemistry of amines [132]. The stoichiometry of the cleavage of, typically, p-toluenesulfonamides in the presence of a suitable proton donor is given by Equation 16.86: O S

H3C

O

R

+2e–, +H+

N

O–

R'

O

R +

S

H3C

NH

(16.86)

R'

The solvents most often used are aqueous dioxane [328], MeOH or aqueous MeOH [260,329,330], MeCN [331], and DMF [332–338]. Lead was found to be a good cathode material for the cleavage of N-tosyl amino acids and peptides [330]. Yields have been reported to be in the range 55–98%. Similarly, benzene-1,4-disulfonic acid diamides may be reductively cleavaged to the amine and the monosulfonamide-sulfinic acid [339]. N-tert-butyl-α-phenylnitrone (BNP) spin trap studies of the reductive cleavage of aromatic sulfonamides under aprotic conditions have been found to be in agreement with a mechanism including the initial formation of a radical anion that subsequently undergoes N–S cleavage to the arenesulfinate ion and an amine radical (Scheme 16.20) [340] that may be further reduced and protonated to the amine. When the reductive cleavage is conducted under aprotic conditions, complications caused by the strongly basic conditions are unavoidable. If the protected amine is primary, the resulting secondary sulfonamide may act as a proton donor in the protonation of the electrogenerated O Ar

S

+e

N

O Ph

R

R N

S

Ar

R'

O

BNP

R'

–

O

R

N R'

C H

R

O Ar

N R'

S

R +

O–

N R'

O N t-Bu

SCHEME 16.20 Reductive cleavage of an aromatic sulfonamide followed by trapping of the resulting aminyl radical.

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601

Cleavages and Deprotections

amide ions, resulting in an overall 1F process and recovery of 50% of the sulfonamide during work-up, Equation 16.87 [334]. O

R +2e– N

S

2H3C

H

O

(16.87) O H3C

S

O

R +

O–

NH + H3C

S

R'

R N–

O

Even when the protected amine is secondary, the resulting tertiary sulfonamide may be deprotonated if the substituents on the nitrogen carry an α-hydrogen atom. Under those conditions, β-elimination with formation of an aldimine or ketimine may occur, Equation 16.88, and the imine may be reducible at the potential required for cleavage of the sulfonamide [335]. O

R

R +2e–

2H3C

S

N

–Ts–

N–

CH R˝

O

CH R˝

R΄

R΄

(16.88)

R

R N

HN Ts–NR–CHR΄R˝ CH R˝ +

–Ts– R΄

C

R˝

R΄

The reduction potentials of aromatic sulfonamides are low [328,332], in 75% dioxane typically –2.3 to –2.4 V vs. NCE [328], and are dependent on the nature of the substituents as expected; electron-donating substituents make the sulfonamide more difficult to reduce, whereas electronwithdrawing substituents make the sulfonamide more easy to reduce. These effects could be rationalized in terms of a Hammett relation [329]. The low potentials needed may cause problems if other reducible functions are present in the protected molecule. For that reason, efforts have been made to find ways to make cleavage occurring more easily, for instance, by indirect electrolysis using a mediator. Indirect electrolysis using radical anions of aromatic hydrocarbons or Ni(acacen) as the mediators has made it possible to determine the formal potentials of a series of tosylamides in DMF and also to estimate the rate of cleavage of the sulfonamide radical anions [333]. The results were similar to those for the tosylates; when the amine was aliphatic, the cleavage of the anion radical was the rate-determining step, whereas the homogeneous electron transfer from the electron donor became rate determining when the amine was aromatic. Other systems for indirect electrolysis of sulfonamides in DMF include anthracene [336] or naphthalene [341] as mediators. Another way to make reduction easier is to use nitrobenzenesulfonyl protected amines [337,338]. The radical anions of N,N-dibutyl-4-nitrobenzenesulfonamide and N-butyl-2-nitrobenzenesulfonamide are persistent in DMF at the time scale of CV, but reduction of the radical anion to the dianion induced a rapid cleavage of the N–S bond with formation of the amine in good yield [337]. Also, the introduction of the N-Boc group was found to make reduction easier by 0.2–0.3 V [342]. Thus, tert-butyl sulfonylcarbamates such as RSO2N(Boc)CH2Ph could be cleavaged in MeCN to BocNHCH2Ph in MeCN.

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Organic Electrochemistry

The electrochemical cleavage in MeCN with Bu4NHSO4 as the electrolyte has been demonstrated to work well for β-amino alcohols protected at nitrogen as the N-benzenesulfonyl derivative and at oxygen as the tert-butyldimethylsilyl derivative or for N-benzoyl substituted derivatives [343]. In N-tosylcarboxamides, competition between the C–N and N–S might be expected. However, during reduction in DMF containing acetic acid, the N–S bond is preferably cleavaged. The yield of the deprotection amine was highest when a mercury cathode was used [344]. Besides being used for the deprotection of amino acids and peptides, electrochemical cleavage has been employed in the synthesis of polyaza and polyaza–polyoxa ligands [345]; the deprotection may be performed by direct as well as by indirect electrolysis. 2. Sulfenamides N-tert-butyl-2-benzothiazolesulfenamide [346,347] and the corresponding morpholino derivative [348] are vulcanization accelerators. Electrochemical reduction at mercury leads to cleavage of the N–S bond and the formation of the substituted amine via the prior formation of a mercury mercaptide [346]. 3. Sulfimides (Sulfilimines) Sulfimides of the general structure shown in Scheme 16.21 undergo reductive cleavage, predominantly of the N–S(IV) bond to give the sulfonamide and the sulfide [349]. A minor pathway includes cleavage of the N–S(VI) bond. A variant of this scheme, presumably including disproportionation of the initially formed radical anion, was suggested for the case Ar = p-nitrophenyl [350].

B. OXIDATIVE N–S CLEAVAGES 1. Sulfinamides The structures of the products resulting from constant current electrochemical oxidation of N-p-toluenesulfinamides depend on the degree of substitution on the nitrogen atom [351]. Sulfinamides derived from secondary alkylamines and primary arylamines undergo oxidative N–S cleavage in MeOH with the formation of the amines and the methyl sulfinate (Scheme 16.22, top), whereas sulfonamides derived from primary alkylamines are oxidized to the corresponding sulfonamides in good yields (Scheme 16.22, bottom). Ar

SO2 N

SolvH Ar –Solv

SPh2

SO2 N

Ar

Ar

SO2 NH2

Reductive cleavage of a sulfimide (sulfilimine). R = alkyl, R΄= alkyl R = aryl, R΄= H

O

R N

S

R'

p-Tol

N

H +

R' Pt cathode RVC anode

S MeO

O

R N R΄

Oxidative cleavage of p-toluenesulfinamides.

© 2016 by Taylor & Francis Group, LLC

O

R

–2e–, MeOH

R = alkyl, R΄= H

SCHEME 16.22

SPh2

SO2 NH + Ph2S +e–, +H+

SCHEME 16.21

–

+e–

S

O p-Tol

p-Tol

603

Cleavages and Deprotections

2. Sulfenamides The radical cations of sulfenamides generally react only slowly [352,353]. However, results from ESR spectroscopy have shown that the radical cations generated by electrochemical oxidation of 4′-unsubstituted N-alkyl-2-nitrobenzenesulfenanilides are indeed due to the dimers that upon further oxidation undergo N–S cleavage and the formation of RS+ [352].

XIV. NITROgEN–HALOgEN bONDS A.

REDUCTIVE N–Hal CLEAVAGES

A study of the reductive cleavage of the N–Hal bonds in saccharin derived N-halosultams and their 4-nitro-substituted analogues (Scheme 16.23) in MeCN demonstrated that the two main factors determining the occurrence of a stepwise or a concerted mechanism are the energy of the π* orbital that accommodates the incoming electron and the strength of the bond to be broken; the weaker the bond, the more chance for the concerted mechanism to overrun the stepwise mechanism [354]. Thus, a concerted cleavage was observed for R = H, Hal = F and R = NO2, Hal = Cl whereas a stepwise mechanism was observed for R = NO2, Hal = F (see also Chapter 14).

XV. PHOSPHORUS–PHOSPHORUS bONDS A. REDUCTIVE P–P CLEAVAGES The polarographic reduction of diphosphines and cyclic polyphosphines in DMF proceeds as irreversible 2F processes resulting in the cleavage of the P–P bond as illustrated by Equation 16.89 for 1,1,2,2-tetraphenyldiphosphine [355]. The mono- and disulfides react similarly: Ph

Ph

Ph +2e–

P Ph

XVI. A.

2

P

P–

(16.89)

Ph

Ph

PHOSPHORUS–OXygEN bONDS REDUCTIVE P–O CLEAVAGES

(Diethylamino)(phenoxy)diphenylphosphonium ions are reduced in aqueous solution with cleavage of the P–O bond and the formation of N,N-diethyl-1,1-diphenylphosphinamine [115]. Triphenylphosphine oxide is a well-known by-product from the Wittig reaction and the possibility of regenerating triphenylphosphine by electrochemical reduction of the oxide has received special attention. Usually, direct electrochemical reduction results in C–P bond cleavage and is for that reason not useful although it has been reported that deoxygenation may be accomplished by reduction in a cell containing a zinc cathode and an aluminum anode [356]. However, the oxide may easily be converted to the dichloride by the in situ reaction with, for instance, oxalyl chloride or trimethylsilyl chloride in MeCN; the resulting dichloride may in turn be reduced to triphenylphosphine in

N

R = H, NO2 Hal = F, Cl, Br

S

R O

SCHEME 16.23

Hal

O

Structure of a saccharin derived N-halosultam.

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Organic Electrochemistry

a high yield in an undivided cell using a platinum cathode and aluminum or zinc anodes, Equation 16.90 [357]. Alternatively, triphenylphosphine may be regenerated via conversion of the oxide to the methylthiotriphenylphosphonium ion followed by reduction (see the following) [358]: Ph P

Ph

O

Ph

(COCl)2

Ph

Ph

Cl

+2e–

P Ph

–2Cl–

Cl

Ph Ph

(16.90)

P Ph

Phenyl diphenylphosphinite undergoes P–O reductive cleavage in fashion similar to that found for halodiphenylphosphines [359] (see Section XIX.A).

B. OXIDATIVE P–O CLEAVAGES The electrochemical oxidation of enol phosphates, phosphites, and phosphinates is accompanied by cleavage of the P–O bond [360]; the reaction proceeds essentially as shown in Equation 16.59 for other esters.

XVII. PHOSPHORUS–SULFUR bONDS A.

REDUCTIVE P–S CLEAVAGES

Ethylthiotriphenylphosphonium ion is reductively cleaved in an irreversible 2F process to triphenylphosphine and ethanethiolate [115]. The reaction has been put to use in the electrochemical regeneration of triphenylphosphine from triphenylphosphine oxide, the Wittig reaction by-product. First, the oxide is converted to the sulfide by reaction with P4S10. Methylation then leads to the methylthiotriphenylphosphonium ion that is finally reduced in MeOH at a mercury cathode to triphenylphosphine and methanethiolate, Equation 16.91 [358]: Ph3P+

O–

Ph3P+

SMe

P4S10

Ph3P+

+2e–

S–

Me2SO4

(16.91)

Ph3P + MeS–

Diphenylarylsulfophosphamides, ArSO2PPh2, are more easily reduced than the corresponding sulfonamides; the radical anions undergo P–S cleavage to the sulfinate anion and the diphenylphosphinyl radical that under the conditions is further reduced to the diphenylphosphine anion, Ph2P– [361]. An in  situ oxygen migration leads to the final products, the arylthiolate and diphenylphosphinate anion, Equation 16.92. O Ar

S O

O

Ph +e– P Ph

Ar

S O

Ph ArSO2– + Ph2P

P Ph

ArS– + Ph2P(O)O–

+e– in situ

(16.92)

ArSO2– + Ph2P–

Triphenylphosphine sulfide and similar compounds are reduced in DMF in two steps; first to the radical anion in a reversible one-electron process [128] and then to the dianion that undergoes P–S cleavage to triphenylphosphine and sulfide ion and C–P cleavage to the mercaptodiphenylphosphine anion and the benzene anion that is further protonated (Scheme 16.24) [127]. Ethylthiodiphenylphosphin is reductively cleaved in a 2F process to the diphenylphosphine anion and thioethanolate [362]. In the absence of a proton donor, the diphenylphosphin anion attacks the starting material in a substitution reaction leading to tetraphenyldiphosphin.

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605

Cleavages and Deprotections +e– Ph3P+

S–

Ph3P

S–

–e– +e– S–

Ph3P

Ph2P

S– + Ph–

+e– Ph3P + SH2–

SCHEME 16.24

Reductive cleavage of triphenylphosphine sulfide.

XVIII. PHOSPHORUS–SELENIUM bONDS A. REDUCTIVE P–Se CLEAVAGES The polarographic reduction of triphenylphosphineselenide in DMF results, in contrast to the reduction of triphenylphosphinesulfide (see Section XVII.A), in cleavage of the P–Se bond in an irreversible 2F process, Equation 16.93 [127]: Ph3P+

XIX.

+2e–

Se–

(16.93)

Ph3P + Se2–

PHOSPHORUS–HALOgEN bONDS

A. REDUCTIVE P–Hal CLEAVAGES The P–Hal bond in halodiphenylphosphines may be cleavaged reductively [359]. A mechanism that involves intermediate diphenylphosphinyl anions was proposed to account for the formation of tetraphenyldiphosphine as the only product (Scheme 16.25). The conversion of the Wittig by-product, triphenylphosphine oxide, to the dichloroderivative and reduction of the latter to regenerate triphenylphosphine [357] has been mentioned in Section XVI.A.

XX. ARSENIC–ARSENIC bONDS A.

REDUCTIVE AS–AS CLEAVAGE

The polarographic reduction of organic compounds containing an As–As bond proceeds analogously to the 2F reductive cleavage observed for diphosphines. Thus, reduction of 1,2-diphenyl-1,2diarsacyclopentane in DMF, for example, results in cleavage of the As–As bond in the five-membered ring as shown in Equation 16.94 [355]. Ph

Ph As

As

+e– Ph2PX Ph2P– + Ph2PX

SCHEME 16.25

+2e–

Ph2PX –

–X–

Ph

Ph

As–

As–

+e– –X

–

Ph2P

Ph2P PPh2

Reductive cleavage of halodiphenylphosphines.

© 2016 by Taylor & Francis Group, LLC

Ph2P–

(16.94)

606

Organic Electrochemistry

H OR O

O

O O

O

O O O H

H O O O

SCHEME 16.26

Examples of cyclic peroxides that undergo reductive O–O cleavage.

XXI. OXygEN–OXygEN bONDS A.

REDUCTIVE O–O CLEAVAGES

Organic compounds with an oxygen–oxygen single bond such as hydroperoxides [363], dialkyl peroxides [364], the acetone peroxide dimer [365], peracids [366], peroxy esters [367], and diacyl peroxides [367] may all be reduced electrochemically. The reaction that has analytical applications includes cleavage of the O–O bond [368] and proceeds for simple hydroperoxides and dialkyl peroxides to the corresponding alcohols. The electron transfer process has been studied in detail for the naturally occurring ascaridole [369] (Scheme 16.26, left) and for the antimalarial artemisinin [370] (Scheme 16.26, middle), and it was found that the reductions followed a concerted dissociative mechanism leading to the distonic radical anions. In contrast, for the G3-factor endoperoxide [371] (Scheme 16.26, R = H) and the methyl ether (R = Me), a transition from a stepwise mechanism to one with a more concerted character was observed. In the presence of a proximal acid group, concerted heavy-atom cleavage and proton and electron transfers may take place [372]. Indirect reduction of di-tert-butyl peroxide and tert-butyl peracetate by aromatic radical anions in DMF generates t-butoxy radicals that abstract a hydrogen atom from DMF. The resulting N,N-dimethylaminocarbonyl radical was observed to undergo coupling with the mediator [364].

XXII. A.

OXygEN–SULFUR bONDS

REDUCTIVE O–S CLEAVAGES

1. Alkoxysulfonium Ions Ethoxysulfonium ions are reduced in MeCN to a mixture of the corresponding sulfide and sulfoxide in an overall 1F process [220]. Spin trapping with α-phenyl-tert-butylnitrone implicated radical intermediates, Equation 16.95; the sulfoxide arises from a subsequent reaction between EtO – and the starting material: Spin trap OEt

Spin adduct +e–

S+ R΄

R˝

R΄

S

R˝ + EtO

(16.95)

+e– EtO–

2. Sulfonates The direct or indirect reductive O–S cleavage of sulfonic acid esters to sulfinate ions and alcohols, Equation 16.96, has been put to use in protection–deprotection chemistry; tosylates [132,332,333,373,374] and nosylates [375] (esters of 4-nitrobenzenesulfonic acid) have featured in particular in this area of chemistry.

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607

Cleavages and Deprotections O H3C

S

OR

+2e–, +H+

SO2– + ROH

H3C

(16.96)

O

Experiments with the esters of l-menthol and l-borneol have demonstrated that the stereochemical identity of the alcohols is preserved during the reduction [373]. It was found that the S–O cleavage takes place at the radical anion stage with a rate constant in the range 104 –108 s−1 when the reaction is carried out under nonaqueous conditions [333]. In the absence of a suitable proton donor, the deprotection is hampered by the nucleophilic attack of the RO – anion on the tosylate with formation of the ether, ROR, as the result, Equation 16.97: Ar

SO2 OR + RO–

Ar

(16.97)

SO3– + ROR

This side reaction can be suppressed by addition of, for instance, acetic acid. In that case, aromatic radical anions cannot be used as catalysts in an indirect process; however, Ni(acacen) is suitable under those conditions. The difference in yield of deprotected alcohol from direct and indirect reduction was insignificant [333]. The esters of benzenedisulfonic acids react similarly to give the mixed sulfonic ester-sulfinic acid and the alcohol [339]. Nosylates are reduced in two steps; the first reduction results in the formation a rather stable radical anion at a potential characteristic of substituted nitrobenzenes [375]. In the second step (–1.2 to –1.7 V vs. SCE), a dianion that cleaves rapidly to ROH and the arylsulfinate ion is formed. Bis(tosyloxy)benzene derivatives may be selectively cleaved; a tosyl group ortho or para to an electron-withdrawing group (e.g., an ester group) is cleaved preferentially to one in the meta position, whereas in derivatives of anisol, the tosyl group meta to the methoxy group is preferentially cleaved [376]. The reduction of alicyclic vinyl triflates in the presence of CO2 results first in cleavage of the O–S bond in an indirect process; in a second step, the enolate ion reacts with CO2, resulting in the formation of a β-keto acid in yields that are often higher than 70%, Equation 16.98 [377]. O

O–

SO2 CF3 +CO2 – –CO2, –CF3SO2–

O +CO2

CO2–

(16.98)

3. Sulfinates The polarographic reduction of methyl benzenesulfinate under aqueous conditions results in the formation of thiophenol and methanol in a 4F O-S cleavage process. Under aprotic conditions a 2F process leads to the benzenesulfenate anion and methoxide ion [378]; it has been suggested that the benzenesulfenate ion is likely to undergo disproportionation to benzenesulfinate and thiophenolate ion. 4. Sulfenates Methyl benzenesulfenate undergoes cleavage of the S–O bond generating thiophenol and methanol in a 2F process when reduced electrochemically in aqueous buffer at pH 7 [379]; methyl 2-nitrobenzenesulfenate reacts similarly when reduced in DMF [380]. A detailed CV study of the reduction of benzyl benzenesulfenate in DMF showed that the S–O cleavage in this case takes place in a stepwise dissociative mechanism [159]. In contrast, the reduction of tert-butyl benzenesulfenate involves cleavage of the C–O bond.

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Organic Electrochemistry

5. Sulfoxides Simple aliphatic sulfoxides are not reduced electrochemically—one of the reasons why DMSO is a popular solvent for electrochemistry. However, a catalytic hydrogen reduction wave is observed during polarography of methyl phenyl sulfoxide in acidic solution [381]. When the reduction is carried out in DMF or MeCN with phenol as the proton donor, conversion to methyl phenyl sulfide is observed [382]. Old reports indicate that diphenylsulfoxide [383] and dibenzylsulfoxide [384] may be reduced to the sulfide at a lead cathode in alcoholic sulfuric acid.

XXIII.

SULFUR–SULFUR bONDS

A. REDUCTIVE S–S CLEAVAGES The reductive cleavage of the sulfur–sulfur bond is a classic reaction in organic electrochemistry and has been observed for di-, tri-, and tetrasulfides [385]. The interest has mainly been focused on disulfides that are easily reduced to the corresponding mercaptans; for instance, a high yield of thioglycolic acid [386] was obtained by reducing dithiodiglycolic acid in 2N sulfuric acid at a lead cathode at a potential of −0.55 V vs. SCE. (This is one of the few examples of a controlled potential electrolysis prior to the invention of the potentiostat.) Similarly, polarographic reduction of alicyclic disulfides gives the corresponding dithiols in reactions that involve adsorption at the mercury electrode [387,388]. Related to these is the reduction of thiolsulfonates to the corresponding sulfinic acids and thiols (Equation 16.99), R = alkyl or Ph [389]: R

SO2 SR

+2e– , +2H+ R

SO2H + RSH

(16.99)

The reduction of aromatic disulfides follows the same pattern and results in the formation of thiophenols [390,391]. The details of both the direct and the indirect reduction of aromatic disulfides have attracted considerable interest over the years [392–397]. From the results of a series of studies of the reduction of para-substituted diphenyl disulfides carried out in DMF by CV and convolution analysis, and indirectly by using electrogenerated radical anions as solution electron donors, it was concluded that the reduction is dissociative leading to cleavage of the S–S bond in a stepwise fashion [395–397]. For the disulfides carrying electron-donating substituents, the inner reorganization energies were particularly high reflecting considerable stretching of the S–S bond upon electron transfer. On the other hand, for electron-withdrawing substituents, the extent of delocalization of the SOMO into the aryl system increases, resulting in a decrease of the reorganization energy for the formation of the radical anions. Reduction of disulfides in the presence of an alkylating or acylating agent gives the S-alkylated or S-acylated products in good yields [398]. If oxygen is present, sulfinic acids may be formed [399].

B. OXIDATIVE S–S CLEAVAGES Diphenyldisulfide may be oxidized at a platinum anode in a mixture of glacial acetic acid and concentrated hydrochloric acid to benzenesulfonic acid [400]. When dimethyl, diphenyl, dibenzyl, and heterocyclic disulfides are oxidized in MeCN in the presence of an alkene, products of acetamidosulphenylation are observed. The reaction has been suggested to proceed as shown in Scheme 16.27 [401]. With cyclic alkenes, a high selectivity for the trans-addition is observed. With terminal alkenes, the terminal sulfides are formed with high regioselectivity. A similar reaction has been used in the electrosynthesis of 2,2-bis(halomethyl)penam derivatives from the related disulfides, Equation 16.100 [402]. It is of interest to notice that a five-membered ring is formed and not a six-membered. Other synthetic aspects of this reaction type have been reviewed [302].

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Cleavages and Deprotections

RS

SR

–2e–

2΄RS+΄ R

RS

SCHEME 16.27

CH3CN

S+

΄RS+΄ +

C+ CH3

N

H2O –H+

RS

NHCOCH3

Mechanism for the oxidation of disulfides in acetonitrile in the presence of an alkene.

R΄CONH

S

H

R΄˝ R΄CONH

H

S

S anodic ox MeOH, HX˝

N CH2X΄

O

N

CH2X΄ CH2X˝

(16.100)

O COOR˝

COOR˝

XXIV. SULFUR–HALOgEN bONDS A.

REDUCTIVE S–Hal CLEAVAGES

1. Sulfonylhalides Sulfonylhalides are in general easy to reduce [328]. The reduction of sulfonyl chlorides has been studied in particular and proceeds with cleavage of the S–Cl bond [403–405]. In the absence of a strong proton donor, sulfinate ion is produced in a 2F process, Equation 16.101 [403]; under sufficiently acidic conditions, the sulfinate ion is protonated and may be further reduced to the thiol in a 6F process, Equation 16.102 [406]: Ar

Ar

SO2Cl

SO2Cl

+2e–

+6e–, +5H+

Ar

Ar

SO2– + Cl–

SH + Cl– + 2H2O

(16.101) (16.102)

When the reduction is carried out in MeCN or DMF in the presence of alkylating agents, mixtures of sulfones and sulfides, resulting from alkylation of the sulfinate and thiolate anions, respectively, may be obtained [406]. 2. Sulfenylchlorides Benzenesulfenyl chloride is reduced in MeCN in a “less than 1F” process to diphenyl disulfide and chloride ions [407]. The mechanism suggested that leads to the disulfide and accounts for the low amount of charge required includes reaction between an intermediate benzenethiolate ion and the benzenesulfenyl chloride starting material producing the disulfide and the reaction between benzenesulfenyl chloride and residual water leading to the thiosulfinate than undergo disproportionation to the disulfide and the thiosulfonate. The details of the electron transfer process have been addressed in detail [408]. For the para-substituted compounds, a dissociative process involving strong interactions between the produced fragments (“sticky” dissociative electron transfer,

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610

Organic Electrochemistry RSe–SeR + Hg

Hg(RSe)2

– + Hg(RSe)2 +2e , +2H 2 RSeH + Hg

SCHEME 16.28

Reductive cleavage of diselenides at mercury electrodes.

see Chapter 14) takes place, whereas the electron transfer process for ortho-substituted compounds is stepwise with through-space S⋯O interactions playing an important role in stabilizing both the neutral molecules and the radical anions.

XXV. SELENIUM–SELENIUM AND TELLURIUM–TELLURIUM bONDS A. REDUCTIVE Se–Se AND Te–Te CLEAVAGES The reductive cleavage of the Se–Se bond and the Te–Te bond proceeds in a fashion analogous to that of the S–S bond that is to the corresponding selenols and tellurols; however, during polarographic studies in aqueous solution [388,391,409–411], the participation of mercury in the reduction process is even more important than for disulfides, illustrated in Scheme 16.28 for the reduction of a diselenide in aqueous solution [388,391,409,411]. A more complex variant of this scheme describes the polarographic reduction in aprotic solvents [410,412].

B. OXIDATIVE Se–Se AND Te–Te CLEAVAGES The oxidative cleavage of diselenides in MeCN takes place essentially as described for disulfides by Scheme 16.27; thus oxidation of, for instance, dibenzyl diselenide in MeCN in the presence of cyclohexene results in the formation of N-(2-(benzylselanyl)cyclohexyl)acetamide [413].

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Organic Electrochemistry Horner, L.; Schmitt, R.-E. Phosphorus Sulfur 1982, 13, 169. Kossai, R.; Emir, B.; Simonet, J.; Mousset, G. J. Electroanal. Chem. 1989, 270, 253. Senboku, H.; Nakahara, K.; Fukuhara, T.; Hara, S. Tetrahedron Lett. 2010, 51, 435. Nyasse, B.; Grehn, L.; Ragnarsson, U.; Maia, H.L.S.; Monteiro, L.S.; Leito, I.; Koppel, I.; Koppel, J. J. Chem. Soc. Perkin Trans. 1 1995, 2025. (a) Coeffard, V.; Thobie-Gautier, C.; Beaudet, I.; Le Grognec, E.; Quintard, J.-P. Eur. J. Org. Chem. 2008, 383; (b) Viaud, P.; Coeffard, V.; Thobie-Gautier, C.; Beaudet, I.; Galland, N.; Quintard, J.-P.; Le Grognec, E. Org. Lett. 2012, 14, 942. Casadei, M.A.; Gessner, A.; Inesi, A.; Jugelt, W.; Moracci, F.M. J. Chem. Soc. Perkin Trans. 1 1992, 2001. Kossai, R.; Simonet, J.; Jeminet, G. Tetrahedron Lett. 1979, 1059. Mukminova, G.R.; Chernykh, G.V.; Novikov, V.T.; Avrutskaya, I.A. Russ. J. Electrochem. 1998, 34, 855. (a) Kim, H.-J.; Jung, K.-H.; Choi, Q.-W.; Kim, I.-K.; Leem, S.Y. J. Korean Chem. Soc. 1991, 35, 673; (b) Horner, L.; Vogt, M. Phosphorous Sulfur Relat. Comp. 1979, 5, 287. Kim, H.-J.; Jung, K.-H.; Choi, Q.-W.; Kim, I.-K.; Leem, S.Y. J. Korean Chem. Soc. 1991, 35, 680. Griggio, L.; Capobianco, G. J. Electroanal. Chem. 1978, 94, 67. Griggio, L. Electrochim. Acta 1982, 27, 749. D’Oca, M.G.M.; Russowsky, D.; Canto, K.; Gressler, T.; Gonçalves, R.S. Org. Lett. 2002, 4, 1763. Sayo, H.; Michida, T.; Hatsumura, H. Chem. Pharm. Bull. 1986, 34, 558. Nelsen, S.F.; Steffek, D.J.; Cunkle, G.T.; Gannett, P.M. J. Am. Chem. Soc. 1982, 104, 6641. Andrieux, C.P.; Differding, E.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 1993, 115, 6592. Matschiner, H.; Krech, F.; Steinert, A. Z. Anorg. Allgem. Chem. 1969, 371, 256. Yanilkin, V.V., Gromakov, V.S.; Nigmadzyanov, F.F. Izv. Akad. Nauk Ser. Khim. 1996, 1320. (a) Yano, T.; Kuroboshi, M.; Tanaka, H. Tetrahedron Lett. 2010, 51, 698; (b) Tanaka, H.; Yano, T.; Kobayashi, K.; Kamenoue, S.; Kuroboshi, M.; Kawakubo, H. Synlett 2011, 582; (c) Kuroboshi, M.; Yano, T.; Kamenoue, S.; Kawakubo, H.; Tanaka, H. Tetrahedron 2011, 67, 5825; (d) Kawakubo, H.; Kuroboshi, M.; Yano, T.; Kobayashi, K.; Kamenoue, S.; Akagi, T.; Tanaka, H. Synthesis 2011, 4091. Lecat, J.-L.; Devaud, M. Tetrahedron Lett. 1987, 28, 5821. Hall, T.J.; Hargis, J.H. J. Org. Chem. 1986, 51, 4185. (a) Schmittel, M.; Steffen, J.-P.; Burghart, A. Chem. Comm. 1996, 2349; (b) Schmittel, M.; Steffen, J.-P.; Burghart, A. Acta Chem. Scand. 1999, 53, 781. Pilard, J.F.; Simonet, J. Tetrahedron Lett. 1997, 38, 3735. Kargin, Yu.M.; Al’fonsov, V.A.; Evtyugin, G. A.; Yakovleva, O. G.; Latypova, V. Z.; Mel’nikov, B. V.; Zamaletdinova, G. U.; Batyeva, E. S.; Pudovik, A. N. Russ. J. Gen. Chem. 1985, 55, 891. (a) Whisman, M.L.; Eccleston, B.H. Anal. Chem. 1958, 30, 1638; (b) Hayano, S.; Shinozuka, N. Bull. Chem. Soc. Jpn. 1970, 43, 2039. (a) Kjær, N.T.; Lund, H. Acta Chem. Scand. 1996, 49, 848; (b) Kjær, N.T.; Lund, H. Electrochim. Acta 1997, 42, 2041. Moryganov, B.N.; Kalinin, A.I.; Mikhotova, L.N. Zhur. Obshchei Khim. 1962, 32, 3476. Parker, W.E.; Witnauer, L.P.; Swern, D. J. Am. Chem. Soc. 1957, 79, 1929. Swern, D.; Silbert, L.S. Anal. Chem. 1963, 35, 880. Hammerich, O. Organic peroxides. In Encyclopedia of Electrochemistry of the Elements, Bard, A.J.; Lund, H., eds.; Dekker: New York, 1978, Vol. XI, p. 316. Donkers, R.L.; Workentin, M.S. Chem. Eur. J. 2001, 7, 4012. Donkers, R.L.; Workentin, M.S. J. Phys. Chem. B 1998, 102, 4061. Najjar, F.; André-Barrès, C.; Baltas, M.; Lacaze-Dufaure, C.; Magri, D.C.; Workentin, M.S.; Tzédakis, T. Chem. Eur. J. 2007, 13, 1174. Constantin, C.; Hajj, V.; Robert, M.; Savéant, J.-M.; Tard, C. PNAS 2011, 108, 8559. Horner, L.; Singer, R.J. Chem. Ber. 1968, 101, 3329. Yousefzadeh, P.; Mann, C.K. J. Org. Chem. 1968, 33, 2716. Pletcher, D.; Stradiotto, N.R. J. Electroanal. Chem. 1985, 186, 211. Civitello, E.R.; Rapoport, H. J. Org. Chem. 1992, 57, 834. Kamekawa, H.; Senboku, H.; Tokuda, M. Tetrahedron Lett. 1998, 39, 1591. Persson, B. J. Electroanal. Chem. 1977, 78, 371. Lindberg, B.J. Acta Chem. Scand. 1970, 24, 1110. Todres, Z.V. Electrokhimiya 1988, 24, 563. Gavioli, G.B.; Davolio, G.; Guidetti, E.S. J. Electroanal. Chem. 1970, 27, 135. Chiorboli, P.; Davolio, G.; Gavioli, G.; Salvaterra, M. Electrochim. Acta 1967, 12, 767.

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Cleavages and Deprotections 383. 384. 385. 386. 387. 388. 389. 390. 391. 392. 393. 394. 395. 396. 397. 398. 399. 400. 401. 402. 403. 404. 405. 406. 407. 408. 409. 410. 411. 412. 413.

619

Fichter, F.; Braun, F. Ber. Dtsch. Chem. Ges. 1914, 47, 1526. Fichter, F.; Sjostedt, P. Ber. Dtsch. Chem. Ges. 1911, 43, 3422. Karchmer, J.H.; Walker, M.T. Anal. Chem. 1954, 26, 271. Larsson, E. Ber. Dtsch. Chem. Ges. 1928, 61B, 1439. (a) Nygard, B.; Schotte, L. Acta Chem. Scand. 1956, 10, 469; (b) Nygard, B. Arkiv Kemi 1967, 28, 75; (c) Nygard, B. Arkiv Kemi 1967, 27, 425. Nygard, B.; Olofsson, J.; Bergson, G. Arkiv Kemi 1967, 28, 41. Barnard, D.; Evans, M.B.; McHiggins, G.; Smith, J.F. Chem. Ind. 1961, 20. Persson, B.; Nygard, B. J. Electroanal. Chem. 1974, 56, 373. Nygard, B. Acta Chem. Scand. 1966, 20, 1710. Kargin, Yu.M.; Latypova, V.Z.; Ustyugova, I.A.; Belyaeva, N.V.; Budnikov, G.K. Zhur. Obshchei Khim. 1979, 49, 2272. Simonet, J., Carriou, M.; Lund, H. Liebigs Ann. Chem. 1981, 1665. Latypova, V.Z.; Yakovleva, O.G.; Ustygova, I.A.; Kargin, Yu. M. Zhur. Obshchei Khim. 1984, 54, 1083. Christensen, T.B.; Daasbjerg, K. Acta Chem. Scand. 1997, 51, 307. Daasbjerg, K.; Jensen, H.; Benassi, R.; Taddei, F.; Antonello, S.; Gennaro, A.; Maran, F. J. Am. Chem. Soc. 1999, 121, 1750. Antonello, S.; Daasbjerg, K.; Jensen, H.; Taddei, F.; Maran, F. J. Am. Chem. Soc. 2003, 125, 14905. Iversen, P.E.; Lund, H. Acta Chem. Scand. 1974, B28, 827. Degrand, C.; Lund, H. Acta Chem. Scand. 1979, B33, 512. Fichter, F.; Wenk, W. Ber. Dtsch. Chem. Ges. 1912, 45, 1373. Bewick, A.; Coe, D.E.; Mellor, J.M.; Owton, W.M. J. Chem. Soc. Perkin Trans. I 1985, 1033. Torii, S.; Tanaka, H.; Sasaoka, M.; Shiroi, T.; Uto, S. Ger Offen DE 3511149, 1985. Mairanovskii, S.G.; Neiman, M.B. Dokl. Acad. Nauk. SSSR 1951, 79, 85. (a) Gourcy, J.G.; Jeminet, G.; Simonet, J. Compt. Rend. Acad. Sci. Ser. C 1973, 227, 1079; (b) Jeminet, G.; Simonet, J.; Gourcy, J.G. Bull. Soc. Chim. Fr. 1974, 1102. Urabe, N.; Yasukochi, K. J. Electrochem. Soc. Jpn. 1959, 27, 201. Takagi, S.; Suzuki, T.; Imaeda, K. J. Pharm. Soc. Japan (Yakugaku Zasshi) 1949, 69, 358. Bontempelli, G.; Magno, F.; Seeber, R.; Mazzocchin, G.A. J. Electroanal. Chem. 1978, 87, 73. (a) Ji, C.; Goddard, J.D.; Houmam, A. J. Am. Chem. Soc. 2004, 126, 8076; (b) Ji, C.; Ahmida, M.; Chahma, M.; Houmam, A. J. Am. Chem. Soc. 2006, 128, 15432. (a) Nygard, B. Arkiv Kemi 1967, 27, 341; (b) Nygard, B. Arkiv Kemi 1967, 27, 405. Nygard, B.; Ludvik, J.; Wendsjo, S. Electrochim. Acta 1996, 41, 1655. Paliani, G.; Cataliotti, M.L. Z. Naturforsch. B 1974, 29, 376. Ludvik, J.; Nygard, B. Electrochim. Acta 1996, 41, 1661. Degrand, C.; Nour, M. J. Electroanal. Chem. 1986, 199, 211.

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17

Reductive Coupling James H.P. Utley, R. Daniel Little, and Merete Folmer Nielsen

CONTENTS I. Introduction ............................................................................................................................. 622 II. Homocoupling by Electrohydrodimerization (EHD) ............................................................. 622 A. Intermolecular Hydrodimerizations ................................................................................ 623 1. General Mechanistic Considerations ..................................................................... 623 2. Common Methods for Mechanistic Analysis......................................................... 624 3. Relationships between Experiments on an Analytical and on a Preparative Scale ...625 4. Stereochemistry ..................................................................................................... 625 5. Slightly Activated Double Bonds ........................................................................... 626 6. Nitriles.................................................................................................................... 628 7. Esters ...................................................................................................................... 631 8. Amides and Oxazolidinones .................................................................................. 636 9. Carboxylic Acids .................................................................................................... 638 10. Miscellaneous Activating Groups .......................................................................... 638 11. Ketones ................................................................................................................... 638 12. Aldehydes ............................................................................................................... 643 B. Intramolecular Hydrodimerizations (Hydrocyclizations) ...............................................644 1. Intramolecular Coupling of Electrophores Linked through the β-Carbons .......... 645 2. Intramolecular Coupling of Electrophores Linked through the EWG Groups......646 3. Intramolecular Coupling of Electrophores Linked through the α-Carbons .......... 647 C. Multiply Activated Alkenes ............................................................................................648 1. α,β-Diactivated Alkenes.........................................................................................648 2. α,α-Diactivated Alkenes ........................................................................................ 652 D. Hydrodimerization of CO2 .............................................................................................. 654 III. Mixed Hydrocoupling of Activated Alkenes .......................................................................... 655 A. Intermolecular Mixed Hydrocouplings ........................................................................... 655 1. Couplings with Acrylonitrile ................................................................................. 656 2. Couplings with Alkyl Acrylate and Crotonate....................................................... 658 3. Other Intermolecular Mixed Hydrocouplings........................................................ 658 B. Intramolecular Mixed Hydrocouplings........................................................................... 659 IV. Reductive Coupling of Alkenes with Other Types of Substrates............................................660 A. Intermolecular Reductive Coupling of Alkenes with Other Types of Substrates ...........660 1. Coupling with CO2 .................................................................................................660 2. Coupling with Aldehydes and Ketones ..................................................................664 3. Coupling with Esters and Anhydrides ...................................................................668 4. Coupling with Organic Halides .............................................................................669 B. Intramolecular Coupling of Alkenes or Alkynes with Other Types of Substrates ......... 676 1. Coupling of Activated Alkenes with Aldehydes and Ketones ............................... 676 2. Coupling of Unactivated Alkenes with Aldehydes and Ketones ........................... 678 3. Coupling of Alkynes or Allenes with Aldehydes and Ketones ............................. 679 4. Coupling of Alkenes with Esters ........................................................................... 679

621

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5. Coupling of Activated Alkenes with Organic Halides .......................................... 681 6. Coupling of Unactivated Alkenes with Organic Halides....................................... 682 7. Coupling of Alkynes with Organic Halides ..........................................................684 V. Reductive Couplings Involving Atoms in an Aromatic Ring ................................................. 685 A. Intermolecular Couplings................................................................................................ 685 1. Coupling between Aromatic Halides ..................................................................... 685 2. Dimerization of Aromatic Compounds Activated with Electron-Withdrawing Substituents ............................................................................................................ 685 3. Coupling of N-Heteroaromatic Compounds........................................................... 690 B. Intramolecular Couplings ............................................................................................... 691 C. Coupling between Aromatic Rings and Other Reagents ................................................ 694 D. Reductive Dimerization of Positively Charged Aromatic Systems ................................ 695 1. N-Alkylpyridinium and Related Systems .............................................................. 695 2. Other Positively Charged Systems ......................................................................... 697 Acknowledgments.......................................................................................................................... 698 References ...................................................................................................................................... 698

I. INTRODUCTION Electroreductive coupling reactions constitute an important class of carbon–carbon-bond-forming reactions involving a multitude of compound types as starting materials. Not all types of coupling reactions are covered in this chapter, and in a number of cases more details can be found in other chapters covering the electrochemistry of specific classes of compounds. In this chapter, the focus is on reductive couplings involving alkenes and other carbon–carbon multiple bonds as a substrate, either in dimerization reactions or in cross-coupling reactions with other compounds. In Section V, carbon–carbon-bond-forming reactions initiated by reduction of aromatic systems are discussed. It is beyond the scope of this chapter to cover the industrial applications of electroreductive coupling reactions. However, the industrial importance of electrohydrodimerizations has initiated an impressive amount of work on this particular type of reductive coupling since the early finding by Baizer that the use of tetraalkylammonium p-toluenesulfonates as supporting electrolyte greatly improved the yield of adiponitrile in the electrochemical reduction of acrylonitrile in partially aqueous media [1]. Further reports from Baizer and coworkers at Monsanto in the 1960s demonstrated the possibility of electroreductive coupling of a great variety of α,β-unsaturated carboxylic acid derivatives (nitriles, esters, and amides) [2–4]. Applications to organic synthesis considered here are mainly concerned with improvements in yields, regio- and stereoselectivity of reductive coupling reactions, in a number of cases by application of transition metal catalysts and templating cations, and on intramolecular reductive couplings. Kinetic and mechanistic studies on reductive couplings are abundant, particularly those exploring possible mechanisms of electrohydrodimerization and dimerization of aromatic systems. It has not been possible in this chapter to use a common scale for reduction potentials because of the many differences in solvent, supporting electrolytes, added acids, electrode material, etc. In some cases, however, the relative values may be of interest in a mechanistic discussion and unless stated otherwise potentials have been measured versus SCE.

II. HOMOCOUPLINg by ELECTROHyDRODIMERIZATION (EHD) Mainly because of the commercial success of the adiponitrile process, hydrodimerization of alkenes activated by electron-withdrawing substituents (Equation 17.1) has been widely and thoroughly explored. Consequently, this section cannot be exhaustive. Concentration on typical reactions and general conclusions may serve as guidelines for further work. Special emphasis will be

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put on the effect of reaction conditions on the mechanisms, product selectivity, and stereochemistry. Section  II.A.1 deals with the monoactivated alkenes, that is, structures of the type 1, where R1 and R2 are H, alkyl or aryl; Section II.B.1 deals with intramolecular coupling reactions where two identically activated alkenes are linked together within the same molecule. The reactions of alkenes activated by two electron-withdrawing groups either in α,α- or in α,β-positions are treated in Section II.C.1. R1 R2

2e–, 2H+

EWG

2 R1

A.

EWG

R2 * R1 * R2

1

(17.1) EWG

INTERMOLECULAR HYDRODIMERIzATIONS

1. general Mechanistic Considerations The reactions of monoactivated alkenes following one-electron reduction may in most cases be rationalized according to Scheme 17.1. Two major side reactions compete with the coupling reaction: protonation of the radical anion followed by further reduction and protonation leading to the saturated dihydro product, and polymerization induced by the basic dianion formed by coupling of two radical anions. Other, less typical reaction pathways include reaction between a radical anion and a molecule of substrate (Scheme  17.2), dimerization of two radicals formed by protonation of the initial radical anion (Scheme 17.3), or, infrequently, cleavage of the radical anion followed by coupling. However, for radical anions derived from monoactivated alkenes, the pathway in Scheme 17.2 has only been unequivocally established as a major pathway in a few cases in which the final zero-electron product is a cyclobutane, that is, a cycloaddition product. The term zero-electron product is used where reactions are initiated by electron transfer and very little charge is transferred, for example, less than 0.1 F. The pathway in Scheme 17.3 relates mainly to alkenes activated by keto or aldehyde groups for reduction in hydroxylic solvents. Under these conditions, radical anions derived from carbonyl compounds are protonated at oxygen, and the resulting enolic radical, H1•, is more difficult to reduce than the starting compound. Consequently, fast dimerization of the enol radicals may compete with further reduction. For other substrate types, especially in aprotic solvents containing added acids, R2

EWG

R1

R1

EWG

e–

R2 R1

1

EWG

R1 R1

1–

R1 R1 R2

EWG –

H

HB –B–

1H

R2 R2

R2 CHD (cyclic hydrodimer)

1–

–

H 1H

HB –B– R2

R2 – R1

1– –1

EWG – – EWG

1

R1 R1

EWG – EWG n × 1 2 R – R1 R2 EWG Polymerization R2

HB –B–

H –

EWG EWG

R1 R1

R2

R2

EWG – H EWG

HB –B–

R1 R1

R2

EWG H H

R2

EWG LHD (linear hydrodimer)

SCHEME 17.1 Overview of reaction mechanisms open to reduction to radical anion.

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R2 EWG H 1 H R 1H2 Hydrogenation

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R1 R2

EWG

e–

R1 1

R2

EWG

– R1 1–

R2

–

EWG

R1

1

R2

1 R1 R1

R2

EWG –

R2

–1–

EWG

1– –1

EWG

R2 R1 R1

EWG

R2

R1 R1

1

R2

R2

EWG EWG – – EWG

Cycloaddition (catalytic)

Dimers as in Scheme 17.1

SCHEME 17.2 "Zero electron," (99 0 39 29

1.32 1.01 1.66 1.51 1.60

79 >99 43 48 39

a b c

DMF, Et4NBF4, div. cell, Hg-cath., CPE: Ew = Epred + 0.1 V. Overall yield >95%. From [31]. C = 8 mM. C = 40 mM.

rise to increasing values of n and enhanced yields of the hydrogenation product as a result of an enhancement of the rate of protonation of the radical anions [31]. From the limited amount of data in Table 17.1, higher substrate concentrations tend to give smaller coulometric n-values and higher yields of the dimer, and it appears that the more easily the substrate is reduced the more favored is dimerization. Allylbenzene and ring-substituted derivatives, 8, show no reduction waves before the background in DMF [32]. However, reduction at constant current gives an electrogenerated base (EGB), probably OH−, which catalyzes isomerization to the conjugated alkenes, β-methylstyrenes, which are reducible at accessible potentials. Isomerization was complete after the passage of 0.5 F, and further reduction resulted in mixtures of the hydrodimer of the β-methylstyrene and the corresponding hydrogenation product (Table 17.2). The dimerization reaction showed a high degree of stereoselectivity, the (±) isomer being the major isomer [32].

TAbLE 17.2 Effect of Substrate Structure on the Competition between Dimerization and Hydrogenation in the Reduction of Ring-Substituted Allylbenzenes, 8a Ar Phenyl 2-Tolylb 3-Tolylc 4-Tolylb 4-t-Butylphenylb 2-Anisyld 4-Anisyle 4-Cyanophenylb b

a b c d e f

yield of Dimer (%)

(±) Isomer (%)

yield of Hydrogenation Product (%)

64 61 65 6 9 41 3 54

84 87 88 86 90 87

14 13 6 55 57 22 57

f f

DMF, Bu4NClO4, div. cell, Pt-cath., CCE. From [32]. 2 F passed. 2.5 F. 4 F. 3 F. Not determined.

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f

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Organic Electrochemistry

Compounds of the type 9 (4-ylidene-substituted cyclopentadithiophenes) form radical anions that in the presence of water dimerize following the RR mechanism (as determined by LSV) with k dim ≈ 105 M−1 s−1, yielding a stereoisomeric mixture of hydrodimers. In the absence of water, the dimer dianions are relatively stable and may on a voltammetric time scale be reoxidized at a potential 0.7–0.8 V anodic relative to the reduction potential of 9 [33]. Interestingly, the polymers (linked through the thiophene rings) formed by oxidation of 9 may also undergo reductive coupling, thereby forming cross-linked polymer chains. Oxidative polymerization of the hydrodimers, formed by reduction of 9 in the presence of water, results in a polymer with the same properties as the one obtained by reduction of a pre-formed polymer in the presence of water [33]. Reduction of 2-vinylpyridine, 10a, in H 2O (Et4NOTs) and of 4-vinylpyridine, 10b, in DMF/ H2O (MeEt3NOTs) give good yields (69% and 82%, respectively) of the LHDs [34]. Later, experiments in DMF have shown that electrolysis of 10 at a potential corresponding to the foot of the reduction wave gives good yields of the 0 F product trans-1,2-dipyridinylcyclobutane. This results from a radical anion catalyzed cycloaddition reaction [35] (see Scheme 17.2), that is, under conditions where a high ratio of substrate to radical anion is maintained. The change in product with the change in working potential indicates that the rate constant for reaction between radical anion and substrate is smaller but comparable to the rate constant for reaction between two radical anions. It was shown [35] that the cyclobutane was reduced to the LHD compound at the higher potentials, implying that the cyclobutane might be the common reactive intermediate rapidly reduced at the higher potentials. The mechanistic details of the catalyzed cycloaddition reaction are not known. The 0 F reaction of 10 is similar to the 0 F reaction of aryl vinyl sulfones (see [36]). 6.

Nitriles Et CN

CN R2 11

R1 12

CN

a: R1 = H, R2 = Me b: R1 = CH2CH2CN, R2 = H c: R1 = Me, R2 = H

Ph

Et + Bu N

+

N Bu Bu

Bu 13

14

The industrially important hydrodimerization of acrylonitrile, 11, to the LHD, adiponitrile, is carried out in water with high chemical yields despite the fact that aqueous conditions normally favor hydrogenation of the double bond. The process was developed on the basis of the finding [1] that when high concentration of 11 (>10%) was electrolyzed in aqueous solution containing high concentrations of Et4NOTs (>0.5 M) as supporting electrolyte, dimerization was favored over hydrogenation. Since then much effort has gone into the investigation of the effect of surfactants on the product distribution and the reaction mechanism for simple EHDs, especially that of acrylonitrile, 11 [37–52]. The main effect of the surfactants seems to be the displacement of water molecules from the electrode surface. The adsorption of such organic additives to electrode surfaces is dependent on the electrode material, the potential, and the nature of other ions in the solution [40,45] and can be probed by differential capacity measurements. Simple monoactivated alkenes such as 11 are so difficult to reduce ( 95%)d 85%e (±):meso = 5

— — — — — — — Trace — — — — Trace 30%b 45%b 75%b — — — — — — — —

[62] [62] [62] [62] [62] [62] [62] [68,69] [68,69] [68,69] [68,69] [68,69] [69] [69] [69] [69] [60] [60] [60] [60] [60] [60] [60] [60] [16]

A: MeCN, Et4NBr, div. cell, CPE, Hg-cath. B: DMF, Et4NOTs, div. cell, CCE (4 F), Cu-cath., 10°C–15°C. C: MeCN, Et4NOTs, div. cell, CCE (4 F), Cu-cath., 10°C–15°C. D: DMSO, Et4NOTs, div. cell, CCE (4 F), Cu-cath., 10°C–15°C. E: i-PrOH, Et4NOTs, div. cell, CCE (4 F), Cu-cath., 10°C–15°C. F: MeOH, Et4NOTs, div. cell, CCE (4 F), Cu-cath., 10°C–15°C. G: DMF, LiClO4, div. cell, CPE, Hg-cath. H: DMF, Et4NBr, div. cell, CPE, Hg-cath. Stereochemistry not reported. The de could not be determined. This result has been challenged [70]. Yield of the decarboxylated product, 3,4-cyclopentanone.

Electrolysis of aryl cinnamates in DMF gives lower yields of the CHD product, mainly as 3,4-diarylcyclopentanone, than is found for the alkyl cinnamates [16]. Furthermore, coupling is no longer exclusive to the (±)-isomer and there is evidence of base-catalyzed hydrolysis (Table 17.4). Examination of the reduction of 15b in a range of solvents [69] showed a significant change in distribution of the dimer on the CHD and the LHD as the proton donor ability of the solvent was changed (Table 17.4). The more protic the solvent, the higher the yield of the LHD. The effect of even more protic conditions on the product distribution was examined for 15b in MeOH/DMF (1:9) containing different amounts of AcOH. Electrolysis at constant current gave mixtures of the

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hydrogenation product, the LHD ((±) and meso) and the CHD (±). The yields of hydrogenation product increased (from 20 to 60%) while the total yield of dimers decreased with increasing amounts of AcOH. Particularly, the yield of the CHD decreased, from 55 to 2 M) or by the use of micelle

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635

Reductive Coupling

forming surfactants such as Triton X-100 ( 108 M−1s−1) than in aprotic solvents. Ethyl 3,3-dimethylacrylate, 16b, gives the LHD upon reduction in DMF/H 2O (Et4NOTs) [57]. Under similar conditions, ethyl crotonate, 16c, and the corresponding ethyl 3-butenoate give the same LHD in near-identical yields, and the only unconverted monomer was 16e [78]. During the electrolysis, the basicity of the medium is sufficiently high for base-catalyzed isomerization of the two species to take place in a manner analogous to the nitriles. The side products correspond to Michael addition of the conjugate base of the substrate (C-2 or C-4) to starting material [78]. An exception to the high reactivity of radical anions derived from ester-activated alkenes lacking aryl substituents is the radical anion derived from methyl 4-t-butylcyclohex-1-enecarboxylate, 2, which undergoes coupling following the RR mechanism with a second-order rate constant comparable to those found for aryl cinnamates [26] (see Table 17.3). The product isolated after reduction was the LHD (3) formed by diaxial coupling, under stereoelectronic control [26] (see Section II.A.4 and Scheme 17.4). The efficient cyclohydrodimerization of trimethyl aconitate (18)–(19) is of particular interest because it is a biomass, renewable resource, easily prepared from citric acid (17) by fermentation of glucose on a 400,000 tons per annum scale. Controlled potential reduction [79] gave a good yield of the most stable of the 16 possible diastereoisomers (Scheme 17.5). The reaction conditions are not “green” but doubtless, by analogy with cyclohydrodimerization of α,β-unsaturated ketones [80], the reduction could be carried out in aqueous conditions and at a lead cathode. Lactones such as coumarin and its derivatives, 20, undergo efficient hydrodimerization at low pH [81]. The mechanism of the reduction of 4-methylcoumarin, 20b, in DMF, MeOH, and MeOH/ H2O has been examined [75]. In all cases, the mechanism of dimerization was found to be of the RR type, again with a considerable increase in the observed second-order rate constant in going from aprotic to protic solvents (Table 17.3) [75]. Efforts to achieve the diastereoselective hydrodimerization of cinnamates and cinnamate analogs have focused upon exploration of a variety of chiral auxiliaries including (–)-menthyl, (–)-8-phenylmenthyl, and (–)-endo-norborneol esters [70], Evans oxazolidinones [82], and camphor-based frameworks. The camphor-based systems 21a–h and the oxazolidinones E CO2H

OH

CO2Me

CO2H

HO2C CO2H 17

HO2C

CO2H

MeO2C

CO2Me 18

CPE DMF/Et4NBr Hg cathode

a: R = H b: R = Me R 20

SCHEME 17.5 Thermodynamic control in formation of 19.

© 2016 by Taylor & Francis Group, LLC

E 75%

E E

CO2Me CO2Me 5%

O

E

E = CO2Me

MeO2C

O

E

19

636

Organic Electrochemistry

22a–e  proved the most effective at delivering both satisfactory yields and good-to-excellent levels of diastereoselection. a: Ar = Ph b: o-MeOC6H4 c: m-MeOC6H4 Ph d: p-MeOC6H4 Ar e: p-FC6H4 f: 2-naphthyl g: 1-furyl h: 3:4-methylenedioxyphenyl

Ph Ph

O O 21

O

a: R1 = (S)-i-Bu: R2 = H b: R1 = (S)-i-Pr: R2 = H N O c: R1 = (S)-Bn: R2 = H * * d: R1 = (R)-Ph: R2 = H R1 R2 e: R1 = (S)-Me: R2 = (R)-Ph 22 O

Of the camphor-based derivatives 21a,c–e,g,h, the electrohydrodimerization worked well, delivering 52–68% of the (R,R)-dimer with diastereomeric excesses in the range of 87–95%. As indicated [70], electrohydrodimerization of ent-21a afforded the expected enantiomer. Removal of the chiral auxiliary using LAH, followed by Jones oxidation and Fischer esterification, allowed the authors to determine the ee to be 92%. The reactions were conducted at a constant current of 75 mA in an undivided cell, with lead serving as the cathode material and platinum as the anode, Et4NOTs as the supporting electrolyte, and acetonitrile as the solvent. For unknown reasons, 21b and 21f proved unsatisfactory, the former leading to only 3% and the latter to 18% of the hydrodimer, 23. In all cases, the chemistry was accompanied by the formation of 22–77% saturation of the α,βunsaturated unit of the starting material, the amount of 24 varying as a function of Ar (e.g., 22% when starting from 21h and 77% starting with 21b). The diastereochemical preference was deemed to be the consequence of what is clearly a least hindered Si-face approach of two radical anions. Ph

Ph

Ph 75 mA, undivided cell

Ph O Ar O 21a–h

Pb cathode, Pt anode Et4NOTs, CH3CN

Ph

+

O

Ph O Ar

Ar O 23a–h

2

O 24a–h

8. Amides and Oxazolidinones Activation of alkenes by amides and N-alkylamides has not been used as much as activation by ester groups. The amides are more difficult to reduce than the corresponding esters but have been reported to give the LHDs in 40–80% yields [2,4]. Chiral N-trans-cinnamoyl-2-oxazolidinones, 22, give upon reduction in dry MeCN the all-trans CHD in high yields with a diastereoisomeric excess up to 66% [83] (see Table 17.5). The products are hydrolyzed and esterified to give the dimethyl 3,4-diphenyladipate with up to 70% ee (Scheme 17.6). The chiral 2-oxazolidinones could be recovered in >90% yield. Both a constant current and a controlled potential electrolysis (−1.8 V vs. SCE) of oxazolidinone 22a were explored. The constant current runs occurred most efficiently at a value of 0.1 A; when it exceeded 0.1 A, or was smaller, then the amount of reduction product 25 increased at the expense of hydrodimerization [82]. Utilization of either the constant current or the controlled potential protocol led to the same level of diastereoselectivity, viz., an 85:15 preference for formation of the R,S,Rketo ester 26 over the S,R,S form, 27. The influence of supporting electrolyte upon stereoselection was also investigated. In contrast with the results obtained when Et4NX salts were used, the use of lithium perchlorate in acetonitrile led to the deposition of lithium metal on the cathode, while in THF only the uncyclized hydrodimer formed and as a mixture of d,l- and meso-forms. Table 17.5 illustrates the results for five substrates, 22a–e, electrolyzed at controlled current (CCE) or controlled potential (CPE). This is a good example of the simplification of conditions to develop a convenient electrosynthesis.

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637

Reductive Coupling

TAbLE 17.5 Diastereoselectivity for Reduction of Oxazolidones: CPEa and CCEb R1

yield (%) and Ratio: 26 (2R, 3S, 4R):27 (2S, 3R, 4S)

R2

22a, (S)-i-Bu

H

22b, (S)-i-Pr

H

22c, (S)-Bn

H

22d, (R)-Ph

H

22e, (S)-Me

(R)-Ph

a b c

75a (95)c 76 (88) 68 (83) 70 (81) 72 (92)

85:15a (83:17)c 83:17 (81:19) 70:30 (70:30) 30:70 (32:68) 75:25 (76:24)

MeCN, Et4NOTs, undiv. cell, Pb cath., CPE, Ew = −1.8 V. From [83]. In all cases, the amount of the 2 F product (25) was 11) does a rate-determining dimerization of A−• take place. The simple vinyl methyl ketone, 39a, is electroinactive at these pH values since the β-hydroxyketone anion is formed [92]. For 37a, dimerization of A−• only takes place in the presence of a surfactant, Triton X-100 [41]. In the intermediate pH range, initially formed A−• is protonated to B• either in fast pre-equilibrium or in a rate-determining step followed by coupling between two B• or between A−• and B•. The rate constant for dimerization of two neutral radicals is high—for 37a (in the presence of a surfactant), the value of kdim was estimated to be > 108 M−1s−1 [41]. Also for 38a, a value of kdim = 4⋅105 M−1s−1 was determined at pH 11 (in the presence of surfactant) where the dimerization is between two A−• [41]. O R4

Ar2 O Ar1

R2

R1 37

R3

R1

R O

O

R2

R3

38

Ar 40

39

a: Ar1 = Ar2 = Ph

a: R1 – R4 = H

a: R1 = R2 = H, R3 = Me

b: Ar1 = Ph, Ar2 = 9-anthryl

b: R1 = R2 = Me

b: R1 = R2 = R3 = Me

c: Ar1 = 4-MeOC6H4, Ar2 = Ph d: Ar1 = 4-MeOC6H4, Ar2 = 2,4,6-Me3C6H2

R3 = R4 = H c: R1 = R2 = R4 = Me, R3 = H d: R1 – R3 = Me, R4 = H

c: R1 = H, R2 = R3 = t-Bu

a: Ar = Ph, R = Me b: Ar = Ph, R = t-Bu

Structural and steric factors are important in determining the coupling position (β-carbon or carbonyl carbon). Vinyl alkyl ketones, 39, and simple alkyl substituted cycloalkenones such as 38 in general give mixtures of dimeric products arising from β,β′-coupling, carbonyl-carbonyl coupling, or mixed carbonyl-β-coupling. In contrast, styryl alkyl ketones, 40, and styryl aryl ketones, 37a, almost exclusively give “normal” LHDs or CHDs arising from initial β,β′-coupling. The CHDs formed by intramolecular addition often eliminate water to give another α,βunsaturated ketone. For mesityl oxide (4-methyl-3-pentene-2-one), 39b, the influence of the reaction conditions on the product distribution was examined in detail [97]. Mixtures of four dimeric species were observed, resulting from initial 4,4′-coupling, 41b–d, or from mixed 2,4′-coupling, 41a. Pinacol formation (2,2′-coupling) was not observed (Equation 17.5).

O

+

+

O 39b

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41a

OH 41b

O

O

O

Red.

+ O 41c

41d

(17.5)

640

Organic Electrochemistry

Optimization of the experimental conditions for selective formation of 41a–d in an undivided cell was examined [97]. The product of mixed coupling, 41a, was favored (88%) by reduction in AcOH/Ac2O (1:1). Most other conditions (strongly acidic, aprotic, or basic) favored 4,4′-coupling. The dehydrated cyclic compound 41c (62%) and the hydrated form 41b (20%) were formed in strongly acidic conditions (pH 1.1). Product 41b (cis/trans = 1:1) was favored by reduction in aprotic solvents (DMF, Et4NOTs) or by basic conditions. Cyclization was best avoided by reduction at elevated temperature (65°C) in a 1:1:1 (v:v:v) mixture of H2O, MeCN, and THF (HOAc/KOAc); this gave 41d (56%) [97]. Hindrance at the carbonyl group (e.g., 41c) favors β,β′-coupling. The radical anion derived from 41c is stable on a voltammetric time scale in DMF (n-Pr4NClO4) and reacts slowly under preparative conditions giving 30% of the LHD, exclusively as the (±) isomer [20]. When water or Li+ is added, the rate of reaction is increased. Electrolysis in a mixture of DMF and aqueous buffer (pH 9 or pH 13) gave an increased yield of the LHD (55–60%) predominantly as the (±) isomer (95% CHD (±) (see text)

[21] [21] [21]

642

Organic Electrochemistry

was obtained exclusively, whereas with Fe2+ or Co2+ the CHD was the sole product. The stereochemistry of the CHD was assigned by NMR to arise from meso coupling [21], but later studies have shown the stereochemistry to correspond to (±) coupling by application of 1H NOESY NMR and X-ray crystallography [81]. The metal cations are expected to form ion-pairs with the radical anions (an anodic shift of the reduction peak was observed). Thus, a templating effect is a likely explanation of the stereoselectivity, and stabilization of the dimer dianion by ion pairing may prevent polymerization. The influence of the same metal cations on the yields and products of reduction of ring substituted chalcones, 37, in DMF (Bu4NBr) is similar to the effect described earlier for styryl alkyl ketones. In all cases, polymerization was prevented and mixtures of the LHD and the CHD in an overall yield of 70–90% were obtained; in most cases the presence of the metal ions favored the CHD [22]. In MeCN, the radical anion derived from 9-anthryl styryl ketone, 37b, dimerizes (kdim = 105 M−1 s−1) forming a stable dimeric dianion, most likely via coupling in the β-position [103]. On a coulometric scale, the stable dimeric dianion can be reoxidized to the starting material at a potential ca. 1 V anodic relative to the reduction of 37b. In contrast, the related 9,10-anthryl bis(styryl ketone) undergoes two consecutive one-electron reductions (ΔEo′ = 36 mV corresponding to the statistical factor for reduction of two identical, noninteracting electroactive groups) with formation of a monomeric dianion, which is completely stable on a coulometric time scale. The lack of electronic interaction between the two styryl-keto groups in the monomeric dianion as well as the rapid dimerization of 37b−• is probably caused by the styryl groups being rotated almost completely out of plane of the anthracene ring. In a few examples, radical anions derived from enones have been trapped by electrophiles other than a proton, and the coupling reaction proceeds through neutral radicals. For instance, 48 is reduced in DMF (Bu4NClO4) to give a stable solution of the radical anion. Subsequent addition of Ac2O gives O-acylation in a fast process. The resulting neutral radicals dimerize slowly enough for examination by ESR [104]. Since the unpaired electron density in the neutral radical as determined by ESR is similar in positions 3, 4, 6, 7, and 9, the radicals may couple in several ways. Preparative scale electrolysis on a 0.4–0.5 g scale, with subsequent addition of acetic anhydride, gave a mixture of variously acetoxylated products of the hydrocarbon 49 [104]. 3

2 1 O

4

9 8

5 6

7 48

49

Another example of a coupling reaction initiated by reaction of the radical anion with an electrophile is the reductive coupling of substituted 4H-pyran-4-thiones, 50, in the presence of alkyl halides (Scheme 17.8) [105,106]. The neutral radical formed by alkylation at sulfur is R1 S

R

O

R2

RX –X–

Red.

50

R1

O

R1

O

R2

R2

O 52

R1

SR

S– DMF, Et4NClO4

1

R2

O

R2

x2 1

R

O

RS

SR

–RSSR

2

R

R2

O 51

R1

45–90%

SCHEME 17.8 Radical-radical coupling; first formed radical not easily reduced.

© 2016 by Taylor & Francis Group, LLC

643

Reductive Coupling

apparently not reduced at the potential of the electrolysis but undergoes dimerization. If the substrate is methylated prior to reduction, the sulfonium cation is reduced more easily than the neutral substrate to give the same dimeric product. The initially formed dimer, 51, eliminates disulfide in an oxidatively induced process during the electrolysis, yielding the final bipyranilidene, 52 [105]. 12. Aldehydes R2 R1 R2

R2

R1

O

a: R1 = R2 = Me b: R1 = H, R2 = Me

R

R2 R1

O

1

OH

R 2 R1

OH

R1

R2 R1

R1

R2

O

53

O

R1

R2

R2

O 56

55

54

57

In general, reduction of alkenes activated by CHO, for example, 53, gives preferentially products of mixed coupling (1,3′-coupling). The 3,3′-coupling only takes place when R1 = H or R2 = H. Both 3,3′- and 1,3′-coupling products undergo cyclization to the cyclopentane, 54, or the tetrahydrofuran, 55, derivatives. Often these eliminate water to give cyclopentene, 56, or dihydrofuran, 57, derivatives. The pinacol formed by 1,1′-coupling is the only linear dimer formed. Ph O MeOH, NaClO4 Red.

Ph

O

Ph Ratios 45

OH

Ph

O

+ Total 56%

Ph OH +

Ph :

O

Ph + Total 24% OH

OH

(17.8)

Ph

Ph

55

O

10

:

90

A mixture of pinacol and cyclized mixed coupling products, 55a and 57a, was formed upon reduction of 3-methylcrotonaldehyde, 53a, at pH 5 (Table 17.7) [107]. Under these conditions, the intermediates undergoing coupling are expected to be the neutral radicals obtained by TAbLE 17.7 Influence of Structure on Product Structure and yields Obtained by Reduction of α,β-Unsaturated Aldehydes, 53 Substrate R = R = Me (53a) 53a R1 = H, R2 = Me (53b) R1 = H, R2 = Ph R1 = Me, R2 = Me2C=CHCH2CH2 (E,E) R1 = Me, R2 = Me2C=CHCH2CH2C(Me)=CHCH2CH2 1

2

Conditions

Products and yields

References

A B C D A A

55a 46%, pinacol 8% 55 28%, 57 10%, pinacol 4% 54 26%, 55 53%, pinacol 15% 55 total 80% (see Equation 17.8) 55 36%, pinacol 18% 55 33%, pinacol 18%

[108] [107] [108] [109] [108] [108]

A: EtOH/acetate buffer (1:1), pH 5, div. cell, Hg-cath., CPE. B: Acetate buffer, pH 5, div. cell, Hg-cath., CPE. C: Acetate buffer, pH 4.7, div. cell, Hg-cath., CPE. D: MeOH, NaClO4, div. cell, Hg-cath., CPE.

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644

Organic Electrochemistry

TAbLE 17.8 Influence of Structure on Hydrocyclizations of 59aa EWg

Link

Products and yields

COOEt COOEt COOEt COOEt COOEt

–C(Me)2– –(CH2)2– –(CH2)3– –O–(CH2)2–O–

HC 57%, cis/trans ≈1:1 HC 34%, cis and trans, 4 F prod. 20%, LHD 18% HC 52%, cis/trans ≈1:4, LHD 8% HC 42% HC 48% only trans, 4 F prod. 19%

CN

–(CH2)2–

HC 16%, cis and trans, oligomers 54%

a

MeCN/H2O, Et4NOTs, Hg-cath., div. cell, CPE. Yields based on current. From [113].

protonation at oxygen. The effect of increasing size of R 2 in 53 on the dimeric products [108] is shown in Table 17.8. In all cases, 55 resulting from the mixed 1,3′-coupling was the major isomer [108]. Unsymmetrical coupling is even more prominent in the reduction of cinnamaldehyde in MeOH, which gives four stereoisomeric cyclic products (total 80% yield), all derived from mixed coupling (Equation 17.8) [109] and probably under thermodynamic control.

B.

INTRAMOLECULAR HYDRODIMERIzATIONS (HYDROCYCLIzATIONS)

When two identical activated alkene functions are included in the same molecule, then intermolecular coupling has to compete with intramolecular hydrocyclization. In most cases, the intramolecular reaction, which corresponds to an overall two-electron process, takes precedence. Few mechanistic studies of intramolecular couplings have been reported. The main question is whether the coupling takes place at the mono radical anion stage in an RS-type reaction (one unit reduced, the other not reduced) or at the bis(radical anion) stage in an RR type reaction (both units reduced). The last case implies weak electronic interaction between the electrophores in the initial state. The formal kinetics and diagnostic criteria for the different mechanistic pathways for intramolecular hydrodimerizations under steady-state conditions (LSV, RDE, and polarography) have been established [110–112]. For the RS-type mechanisms, it is normally assumed that the cyclized radical anion (or the neutral, cyclic radical formed by subsequent protonation) is more easily reduced than the starting material. For the RR-type mechanisms, ΔEo′ (the difference in reduction potential of the two electroactive groups) is normally assumed to equal the statistical difference expected for two electronically isolated groups in the same molecule. The two activated alkene functions may be linked via the β-carbons, 58a, via the activating group, 58b, or, less commonly, via the α-carbons, 58c. Where the two alkene units are linked by direct bonding between the α- or the β-carbons, the two alkenes form a single, conjugated π-system. R EWG

R

EWG

R EWG EWG

EWG

R 58a

© 2016 by Taylor & Francis Group, LLC

R

EWG

R 58b

58c

645

Reductive Coupling

1. Intramolecular Coupling of Electrophores Linked through the β-Carbons The efficiency of reductive coupling of 59a depends on the size of the resulting ring, 3–6-membered rings in general being favored (Table 17.8) [113]. In the medium of choice, MeCN/water mixtures, the hydrocyclization (HC) products formed from the general structure 59a by coupling in the β-position were mixtures of the cis and trans isomers (Equation 17.9), analogous, respectively, to the bimolecular meso and (±) couplings in the formation of LHDs. In most of the cases [113], hydrogenated products (2 F and 4 F) and oligomers were formed as well. EWG

EWG R

R

EWG

R

2e–, 2H+ EWG

R

(17.9)

+ R

R

EWG

EWG Trans

Cis

58a

COOR´

R

a: R = H, R´ = Et b: R = R´ = Me

R

COOR´ 59

A systematic study of the influence of added water on the reduction of 59a in MeCN revealed that in the presence of 20% water the HC product (≈90%) was obtained as a mixture of cis and trans isomers in a ratio of ≈1:3, but with a charge consumption slightly lower (1.8 F) than the expected 2 F. Decreasing the water content gradually led, with concomitant reduction in charge, to decreasing amounts of HC product and increasing amounts of cyclic (0 F) products arising from base-induced intramolecular Michael addition [114]. With no added water, the HC product (27%) and the Michael addition product (57%) were formed with consumption of 0.4 F. The suggested mechanistic interpretation of these results were that at low water concentrations the intramolecular coupling involves radical addition to the unreduced function (i.e., an RS-type coupling) followed by hydrogen atom abstraction from the solvent, MeCN (the radical, •CH2CN, being consumed in an unidentified, nonreductive chemical reaction). The resulting anion of the HC product abstracts a proton from unreacted substrate that initiates intramolecular Michael addition (see Scheme 17.9). At higher water concentrations, the radical anion is stabilized. This allows formation of the bis(radical anion) and subsequent RR-type coupling. At the same time, the water serves as proton donor, thereby inhibiting the Michael addition [114]. A similar relationship between product distribution and water content was observed for the analogous substrate containing an extra methylene unit in the link [114]. The Michael addition products were slightly more difficult to reduce than the parent compounds [114]. Formation of the trans isomer of the HC product is enhanced in the presence of metal cations and by using a carbon rather than an oxygen acid. For reduction of 59a in MeCN, the trans:cis isomer 59a

–

e– MeCN (dry) COOEt COOEt

–

MeCN COOEt COOEt – CH2CN –

COOEt COOEt

– COOEt COOEt COOEt –

COOEt

SCHEME 17.9 Intramolecular coupling via Michael addition.

© 2016 by Taylor & Francis Group, LLC

59a

59a

– COOEt + COOEt

COOEt

COOEt –

COOEt

COOEt

COOEt

+ COOEt

646

Organic Electrochemistry

ratio changed from 2.6:1 in MeCN/H2O (9:1) using AcOH as proton source to 7.5:1 in MeCN using diethyl malonate as proton source. The selectivity was further enhanced to 14.8:1 when 1.3 equivalents of CeCl3 was added with a total 73% yield of the HC product [19]. As for the bimolecular EHD reactions, the (±)-coupling is preferred in the presence of templating cations. Intramolecular coupling by the RS mechanism would be expected for reduction of 60. Since the two halves of the molecule interact electronically through the aromatic ring, the first and the second electron transfers are separated by ≈300 mV [114]. Consequently, only the monoreduced substrate can participate in coupling following the first electron transfer. However, only intermolecular coupling products (59%) were found upon reduction [115]. The fully hydrogenated monomer (28%) and the half-hydrogenated monomer (7%) were also formed with an overall consumption of 1.1 F. In the light of the mechanism for alkyl cinnamate reductions, the intermolecular coupling will be of the RR type, and the initial dimer, 61, can then under the conditions of the electrolysis undergo one or two Michael additions to yield 62 and 63 (62%). The products 62 and 63 were shown by X-ray crystallography to result from initial (±) coupling, whereas 61 resulted from meso coupling [115]. COOMe H

CH2COOMe

DMF, 4% H2O

+ MeOOCCH2

MeOOCCH2 H

1.1 F COOMe

meso 60

61

COOMe H +

H

MeOOC

MeOOC

CH2COOMe

CH2COOMe COOMe H

COOMe

H MeOOC CH2COOMe

(±)

(±)

62

63

2. Intramolecular Coupling of Electrophores Linked through the EWg groups Intramolecular reductive coupling of the substrates 64a–c, where (chiral) diols have been used to link two cinnamoyl units through the activating group, has been tried [83]. The yields of HC, intramolecular cyclized products, 66, are normally low (105 2.0⋅105 37 46h

[58] [72] [120] [7] [7]

ESR DPSC/DPSCC DPSC

30 ± 5 44 5.5⋅103 i

[72] [73] [122]

DPSC

2.22⋅105 i

[122]

CV RRDE CV ESR SECM DPSC/DPSCC DPSC/DPSCC RRDE RRDE RRDE DCV

0.1 1.1⋅102 1.6⋅102 (1.60 ± 0.26)⋅102 1.7⋅102 1.2⋅102 25 9.1⋅104 1.9⋅105 6.9⋅104 7.3⋅105

[54] [58] [58] [72] [120] [73] [73] [123] [123] [123] [124]

DCV

5.6⋅105

[124]

CPSV

1.3⋅107

[125]

DCV

6.6⋅105

[124]

DCV

2.2⋅106

[124]

DCV

>107

[124]

n = 0.44. n = 0.6, increases to 1 when Li+ is added. n = 0.5, even when i-PrOH or HOAc are added as proton donors. n = 0.62. Rate constant for isomerization of 80a−• to 79a−• equal to 6.0 s−1. Rate constant for isomerization of 80b−• to 79b−• equal to 2.2 s−1. Rate constant for isomerization of 80c−• to 79c−• equal to 6.7 s−1. Ea = 4.6 kcal mol−1. ΔH≠ = −3.5 kcal mol−1, ΔS≠ = −50 cal K−1 mol−1 when CH2O in the range 70–278 mM.

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kdim(RR)/M−1 s−1 5

651

Reductive Coupling

Large-scale EHD of alkyl fumarates and maleates has been studied, since the tetracarboxylate product has industrial interest [127]. In MeOH, the undesired hydrolysis of 80b is prevented during electrolysis. The use of sodium or lithium acetate as supporting electrolyte rather than R4N+ salts enhances the formation of dimer compared to hydrogenation product[127]. The best preparative results (>80% LHD, ≈10% hydrogenation product, and ≈5% resulting from addition of MeOH to the double bond) were obtained using a graphite felt cathode in an undivided flow cell [127]. There are significant differences in the solubility of 80b and of 79b in MeOH: whereas 80b is miscible with MeOH up to 50% w:v, 79b is only soluble in 3% w:v. The rate constant for dimerization was found by application of a RRDE to be 107 M−1 s−1 for both radical anions under these conditions [128]. At the same time, isomerization of 80b to 79b took place at the radical anion stage. From FT-IR measurements in MeOH (NaOAc), it was found that the amount of 79b present in the solution increased during the electrolysis, and it was estimated that 30% of the 80b initially reduced was converted into 79b [128]. When substituted N-ethyl maleimides, 81, were reduced in aqueous buffers at neutral pH, two of the substrates, 81a and 81b, gave predominantly the LHDs, whereas the third substrate, 81c, only gave the hydrogenation product (Table 17.11). Under acidic conditions, all three substrates were reduced to the hydrogenated monomer [130]. Hydrodimerization of unsymmetrically α,β-diactivated alkenes may in principle give rise to three types of coupling products: the α,α′-LHD formed by coupling between the two α-carbons, the β,β′LHD formed by coupling between the two β-carbons, and the α,β′-LHD formed by coupling between the α-carbon in one molecule and the β-carbon in the other. Examples are given in Table 17.11.

TAbLE 17.11 Hydrodimerization of α,β-Diactivated Alkenes Substrate NC–CH=CH–CN (78) E-EtOOC–CH=CH–COOEt (79a) Z-EtOOC–CH=CH–COOEt (80a) Z-BuOOC–CH=CH–COOBu E-[Me(CH2)3CH(Et)CH2OOCCH=]2 Z-[Me(CH2)3CH(Et)CH2OOCCH=]2 E-Ph(O)C–CH=CH–C(O)Ph 81a (R1 = R2 = H) E½ = −0.77 81b (R1 = Me, R2 = H) E½ = −0.90 81c (R1 = R2 = Me) E½ = −1.02 NC–CβH=CαH–C(O)Ph NC–CβH=CαH–COOEt EtOOC–CβH=CαH–C(O)Ph Me2NOC–CβH=CαH–COOEt a

b c d

e

Conditionsa

Products and yields

References

A B B C D D B E E E Fe B F B

LHDb LHD 83%c LHD 94%, 2 F prod. 5%d LHD 70%c LHD 80%c LHD 66%c LHD 44%b,c LHD 99%, 2 F prod 1%, n = 1.12 LHD 79%, 2 F prod. 21%, n = 1.31 LHD 0%, 2 F 100%, n = 1.73 β,β′-LHD 55%b,c β,β′-LHD 80%, α,β′-LHD 4%c β,β′-LHD 69%,a α,β′-LHD 21%c β,β′-LHD 63%, α,β′-LHD 30%c

[2] [4] [129] [2] [2] [2] [4] [130] [130] [130] [4] [4] [4] [4]

A: MeCN/H2O, Et4NOTs. B: DMF/H2O, Et4NOTs, Hg-cath., div. cell, CCE. C: H2O, MeBu3NOTs, Hg-cath., div. cell, CCE. D: EtOH/H2O, MeBu3NOTs, Hg-cath., div. cell, CCE. E: EtOH/buff. (1:3), KNO3, pH 6, Hg-cath., div. cell, CPE. F: acrylonitrile/H2O, Et4NOTs, Hg-cath., div. cell, CCE. Mixture of meso and (±). Based on current passed. Based on substrate consumed. Two stereoisomers of the LHD were formed in the ratio 93:7 but not assigned. Acid added during electrolysis to keep solution neutral.

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2. α,α-Diactivated Alkenes A significant difference is expected, between α,β- and α,α-diactivated alkenes, since α,αdiactivation will tend to localize the unpaired spin density at the β-carbon, thereby facilitating dimerization. Reduction of 3-carbethoxy- and 4-carbethoxy-coumarin, 82 (α,α-diactivated) and 83 (α,β-diactivated), in DMF and in aqueous/ethanolic buffers [131] shows distinct differences in the competition between dimerization and hydrogenation. In hydroxylic media, 83 leads to the hydrogenation product under acidic conditions and to product mixtures (1–2 F) under basic conditions. In DMF, 83− • is stable on a voltammetric time scale. In contrast, reduction of 82 is a 1 F process in both hydroxylic media (independently of pH) and in DMF, the product being a diastereoisomeric mixture of LHDs [131]. O

O

O

O

COOEt 82

83

COOEt

a: R1 = R2 = Me b: R1, R2 = –(CH(Me)CH2CH(Me)CH2CH(Me))– a: Ar = Ph, R = H 1 2 Ar CN b: Ar = 4-tolyl, R = H R R1 CN c: R , R = –(CH2CH2CH(t-Bu)CH2CH2)– c: Ar = 4-anisyl, R = H d: R1, R2 = –(CH2)5– CN e: R1, R2 = –(CH ) – CN d: Ar = 4-F–C6H4, R = H R2 R 24 84 e: Ar = 4-CN–C6H4, R = H 85 f: R1 = Me, R2 = Et f: Ar = Ph, R = Me g: R1 = Me, R2 = i-Pr h: R1 = R2 = Et i: R1 = Me, R2 =t-Bu

O O O

a: R = 4-F–C6H4 b: R = i-Pr

O 86

The presence of two electron-withdrawing groups at the α-carbon reduces the basicity of the radical anions considerably, and, for example, the radical anions derived from 84a [134,136] and 86a [132] are not protonated by AcOH in DMF and MeCN, respectively. Also, alkylidenemalonitriles, 85, dimerize in competition with hydrogenation in the presence of AcOH (Table 17.12). Kinetic data for the dimerization of the radical anions of substituted benzylidene malonitriles, 84, were in agreement with the simple RR mechanism, and the rate of reaction was not influenced by the concentration of water (Table 17.10) [124]. Compound 84b was studied over a very broad concentration range (2–100 mM) by chronopotentiometry [133] in order to investigate possible mechanistic changes in going from the low concentrations normally applied for analytical work to the high concentrations suited for preparative work. When the water content in the DMF solution was high relative to the substrate concentration (10% water), the dimerization followed the RR mechanism independently of substrate concentration. In dry DMF, the data indicated that although the mechanism was unchanged, changes in the (specific) solvation of the radical anions with increasing substrate concentration affected the rate of the dimerization step [133]. The same mechanistic conclusion was drawn in a study of 84b in MeCN with various amounts of water by CPSV, where the rate constant for dimerization was considerably higher than the one measured in DMF [125] (see Table 17.10). On a preparative scale, different ratios of cis- and trans-isomers of the CHD, 2-amino-4,5diaryl-2-cyclopentene-1,1,3-tricarbonitrile, formed upon reduction of benzylidenemalonitrile, 84a, under apparently identical conditions (DMF, 1 M AcOH) have been reported: cis:trans = 1:1 [134] and exclusively trans [135]. In a later systematic study of the reduction of 84a in DMF (1 M AcOH), the cis:trans ratio was shown to change within the range 1:1–0:1 as a function of substrate concentration, nature of supporting electrolyte, and the degree of conversion. Only the trans isomer was formed when an electron transfer catalyst was used [136]. Small cations and low substrate concentrations favored formation of the cis-isomer, and progressively more cis was formed as the electrolysis progressed. The authors invoke a competition between the

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Reductive Coupling

TAbLE 17.12 Hydrodimerization of α,α-Diactivated Alkenes Substrate 84a 84a 84f 84f 85a 85b 85b 85c 85d 85e 87a 87e 87a 87b 87c 87d 87e a b

Conditions

Products and yields

References

DMF, Bu4NI, 1 M AcOH, C-cath., div. cell, CPE MeCN, LiClO4, 1 M AcOH, Hg-cath., div. cell, CPE DMF, Bu4NI, 1 M AcOH, C-cath., div. cell, CPE MeCN, LiClO4, 1 M AcOH, Hg-cath., div. cell, CPE MeCN, LiClO4, 1 M AcOH, Hg-cath., div. cell, CPE DMF, LiClO4, 1 M AcOH, Hg-cath., div. cell, CPE MeCN, LiClO4, 1 M AcOH, Hg-cath., div. cell, CPE DMF, Bu4NI, 1 M AcOH, Hg-cath., div. cell, CPE DMF, Bu4NI, 1 M AcOH, Hg-cath., div. cell, CPE DMF, Bu4NI, 1 M AcOH, Hg-cath., div. cell, CPE DMF/H2O, Et4NOTs, Hg-cath., div. cell, CCE DMF/H2O, Et4NOTs, Hg-cath., div. cell, CCE MeOH, NaI, GC-cath., undiv. cell, CCE, 60°C MeOH, NaI, GC-cath., undiv. cell, CCE, 60°C MeOH, NaI, GC-cath., undiv. cell, CCE, 60°C MeOH, NaI, GC-cath., undiv. cell, CCE, 60°C MeOH, NaI, GC-cath., undiv. cell, CCE, 60°C

CHD 95% (±):meso = 1:1a CHD 85%b CHD 90% (±):meso = 3:2a CHD 93%b LHD 87%b CHD 60%b CHD 75%b CHD 70%,b 2 F prod. 15% CHD 70%,b 2 F prod. 20% CHD 64%,b 2 F prod. 15% LHD 90%a LHD 45%a 88a 73% trans/cis = 1.35 88b 69% trans/cis = 1.30 88c 58% trans/cis = 1.23 88d 52% trans/cis = 1.17 88e 56% trans/cis = 3.00

[134] [138] [134] [138] [138] [138] [138] [138] [138] [138] [2] [2] [137] [137] [137] [137] [137]

Based on current passed. Stereochemistry not reported.

RR and the RS mechanism as a possible explanation although not all the experimental data are in accord with the suggestion [136]. Reduction of alkylidene malonates, 87, in MeOH in an undivided cell using alkali metal halides as supporting electrolytes results in the unusual formation of 3,4-disubstituted 1,1,2,2-cyclobutanetetracarboxylates, 88 [137]. Cyclobutane formation requires 4–7 F and is not a radical anion catalyzed cycloaddition. The process was explained by the mechanism in Scheme 17.11, where the cyclization takes place by chemical oxidation of the dianion of the LHD by anodically formed halogen. The LHD, 89, was formed initially as determined by interruption of the electrolysis after the passage of 2 F. The yield of 88 decreased when the size of R increased (Table 17.12), and the trans isomer was formed in slight excess. Electrolysis using NaClO 4 as supporting electrolyte gave only the LHD, 89, with a (±):meso ratio close to one. Hydrogenation, addition of MeOH to the double bond, and oxidative rearrangement of the substrate were found to be side reactions [137].

2

R

COOMe

MeOH, Nal CCE, Undiv. cell

COOMe 87

COOMe R

COOMe COOMe

R

COOMe 89

SCHEME 17.11

H+

R R

COOMe – COOMe COOMe – l2 COOMe –2l–

R R

COOMe COOMe COOMe COOMe 88

a: R = Me b: R = Et c: R = Pr d: R = heptyl e: R = Ph

Reversible dimerization of radical anions generated from α,α-diactivated alkenes.

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The derivatives of Meldrum’s acid, 86a,b, undergo reductive dimerization in MeCN, and the mechanism has been studied [132,139]. An RS mechanism with rate-determining electron transfer has been invoked either as the only mechanism (for 86b) or as one of two parallel reaction pathways (for 86a) in order to explain the apparent reaction orders found by DCV or LSV in the presence of AcOH [132]. However, AcOH (pK(AcOH) = 12.3 in DMSO [140]) is not likely to be strong enough to protonate the dimer dianion, since the basicity of the dimer dianion is expected to be close to that of the conjugate base of Meldrum’s acid (pK(Meldrum’s acid) = 7.3 in DMSO [141]). Consequently, the dimerization may be reversible, and this, in turn, may lead to anomalous apparent reaction orders although the coupling is of the RR type [71].

D.

HYDRODIMERIzATION OF CO2

Carbon–oxygen and carbon–nitrogen multiple bond functionality may undergo electrochemical reduction resulting in carbon–carbon bond formation. This category includes aldehydes, ketones, azomethines, esters, and other carboxylic derivatives all of which are covered in other chapters. Carbon dioxide, O=C=O, is difficult to place within the categories referred to. However, its importance as a feedstock for C-1 chemistry, together with the environmental benefit to be gained by discovering ways to convert it into useful products, suggests that attempts at electrochemical conversion should be outlined. Direct reduction of CO2 takes place at rather low potentials (0.4 V, and direct formation of the dianion of 70b takes place before the reduction of 11. Even in the presence of ≈50-fold excess of 11, the perturbation of the RRDE and CV response of 79b is modest at the potential where 79b−• is formed, giving rise only to a slight increase in the rate of disappearance of 79b−• [153]. The voltammogram of diethyl maleate, 80a, is, unlike that of 79b, significantly altered in the presence of a 50-fold excess of 11. The reduction potentials of the maleates are shifted approximately 0.3 V cathodic relative to the reduction potentials of the fumarates (see Section II.C.1) [153], and electron transfer from 80a −• to 11 is therefore expected to be faster than from 79b −•. However, dimerization of 80a−• is much faster than of 79b −• (see Section II.C.1), and since the presence of 11 affects the voltammogram of 80a, reaction between 80a−• and11 must also be fast. Co-electrolysis of equal amounts of 11 and α-methyleneglutaronitrile, 12b, (11 and 12b are reduced at about the same potential) gives a mixture of the three products (the two LHDs and the MHC) in relative amounts corresponding to random coupling of the radical anions [56]. In this case, the two components are structurally similar, suggesting that the individual dimerization rate constants may also be similar. The alkylidenemalonitriles, 85, which are all reduced around −1.7 V, gave modest amounts of the MHC when co-electrolyzed with a twofold excess of 11 (Table 17.13), the yield decreasing with increasing size of R2.

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Reductive Coupling

TAbLE 17.13 Mixed Hydrocouplings with Acrylonitrile, 11, and Ethyl Acrylate, 16d Substrate

Csubstrate/CX

X = 11 Hydrocarbons Ph-CH=CH2 (≈ −2.4 V)

3:1a

Ph2C=CH2 (≈ −2.2 V)

3:1a

Conditions

References

MHC 48%,b LHD of 11, 2 F of Ph-CH=CH2 (major)

[151]

MHC (major), LHD of 11, LHD of Ph2C=CH2, 2 F of 11

[151]

MHC, LHD of 11, LHD of 12b (statistical distribution) MHC 54%c

[56] [163]

MHC 57%d,e

[164]

MHC: 85a 24%, 85f 19%, 85g 13%, 85h 14%, 85i 6%, 85e 21%c

[165]

MHC 50%, LHD of 16a 2.5%, LHD of 11 15.5%c

[166]

MHC 60%,d LHD of 11

[167]

MHC, LHD of 11 major, LHD of CH2=CHCOOEt minor

[164]

MHC 30%,d LHD of 11

[164]

MHC

[4]

MHC, LHD of 80a minor, LHD of 11, 2 F of 80a major MHC 11%, LHD of 80a 80%d

[152]

1:2

DMF/H2O, Et4NOTs, Ew = −1.5 Vg DMF/H2O, Et4NOSO2Ph, Ew = −1.3 to −1.4 Vg

f

H2O, Et4NOTs

Me2C=CHCOMe (39b)

1:10

PhCH=CHCOMe (40a)

1:8

Et4NOTs, H2O, Ew = −1.65 to −1.7 V Et4NOTs, Ew = −1.43 to −1.5 Vg

MHC minor, CHD of 39a, LHD of 11 MHC major

Difunctional trans-NCCH=CHCOOEt

1:>20

Nitriles CH2=C(CN)CH2CH2CN (12b) NCCH=CHCH2CH2CN (≈ −1.9 V) CH2=CHCH=CHCN (−1.5 V) 85a,e–i (≈ −1.7 V)

1:2

Esters CH2=CHCOOMe (16a)

1.8:1

1:1 1:1 1:6

CH2=CHCOOEt (16d) (≈ −1.8 V) 16d (≈ −1.8 V)

1:5

CH2=(COOEt)(NHCOMe)

1:10

trans-EtOOCCH=CHCOOEt (79a) (−1.54 V) cis-EtOOCCH=CHCOOEt (80a) 80a (≈ −1.4 V)

f

Ketones CH2=CHCOMe (39a)

1:1

1:>10h

DMF, 2.5% H2O, Et4NOTs, div. cell, Ew = −2.45 V DMF/H2O, Et4NOTs, div. cell, Ew = −2.30 V

Products and yields

DMF/H2O, MeEt3NOTs H2O, K2HPO4, div.cell, CCE H2O, Et4NOTs, Ew = −1.71 V H2O, K2HPO4, pH 5–6, div. cell, Pb cath. H2O, K2HPO4, pH 6–8, undiv. cell, graphite cath., CCE H2O, Et4NOTs, Ew = −1.8 V DMF/H2O, MeEt3NOTs, Ew = −1.83 to −1.85 V H2O, Et4NOTs, Ew = −1.75 V H2O, Et4NOTsg

DMF/H2O, Et4NOTs, Ew = −1.5 Vg

[164]

[152] [102]

MHC small amount, LHD of 40a major

[102]

MHC, LHD of NCCH=CHCOOEt major, 2 F of NCCH=CHCOOEt

[152]

(Continued)

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Organic Electrochemistry

TAbLE 17.13 (Continued) Mixed Hydrocouplings with Acrylonitrile, 11, and Ethyl Acrylate, 16d Substrate

Csubstrate/CX

Conditions

Products and yields

References

trans-NCCH=CHCOOEt (−1.45 V) transMe2NCOCH=CHCOOEt (−1.76 V) X = 16d 9-benzylidenfluorene (6)

f

H2O, Et4NOTsg

MHC trace

[4]

f

H2O, Et4NOTsg

MHC 50%,d,i LHD of Me2NCOCH=CHCOOEt, 2 F of Me2NCOCH=CHCOOEt

[4]

1:10

MHC 10n

DMF, H2O, Me3BuNOTs, Ew = −1.5 to −1.0 Vj H2O, K2HPO4, div. cell, CCE MeCN, H2O, Et4NOTs, Ew = −1.67 to −1.70 V H2O, K2HPO4, pH 5–6, div. cell, Pb cath. DMF, H2O, Et4NOTs, Ew = −1.5 Vj

MHC, LHD of 80a minor, 2 F of 80a major, 2 F of 16d

[152]

1:8

2 F indicates the 2 F reduction product. a 11 added gradually. b 5-phenylpentannitrile and 4-phenylpentannitrile (15:1), yield based on 11. c Yield based on current. d Based on substrate. e 1,6-dicyanohexane isolated after catalytic hydrogenation, that is, product derived from coupling in δ-position. f 11 as solvent, substrate added slowly during electrolysis. g AcOH added during electrolysis to keep medium acidity constant. h Substrate added slowly during electrolysis. i Coupling α to COOEt/coupling α to CONMe (1:2). 2 j AcOH added during electrolysis to keep medium acidity constant. k Based on substrate. l Methyl acrylate was used instead of ethyl acrylate. m The product derived from coupling in δ-position was isolated after catalytic hydrogenation (in small amounts). n Substrate added slowly during electrolysis.

2. Couplings with Alkyl Acrylate and Crotonate Simple alkyl acrylates have been used as partner in mixed hydrocoupling reactions in a number of cases (Table 17.13). Like mixed couplings with 11, LHD formation and hydrogenation compete with formation of the MHC of alkyl acrylates. Co-electrolysis of ethyl 3,4-dimethoxycinnamate (E½ = −1.94 V) with ethyl crotonate (E½ = −2.37 V) in a molar ratio of 1:7.2 in MeCN at Ew = −2.03 V gave none of the MHC but only the CHD of the cinnamate. The only product derived from the crotonate was diethyl 2-ethylidene-3-methylglutarate (49%) formed by action of an EGB, most likely the dimer dianion of the cinnamate [62]. 3. Other Intermolecular Mixed Hydrocouplings Reduction of diethyl fumarate, 79a, and the simple enone 39a (1:10) in MeCN/water (Et4NOTs) at a potential between the two reduction potentials (ΔE ≈ 0.2 V) gave the MHC as the major product [102].

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Reductive Coupling

Reduction of dibutyl maleate 80c, and 2-vinylpyridine, 10a (1:5), in DMF/water (Et4NOTs) at a potential where only the maleate is reduced (ΔE ≈ 0.15 V) leads to a mixture of the two LHDs and the MHC. Analogous results were found for co-electrolysis of 4-vinylpyridine, 10b, with an excess of the simple enone, 39a (ΔE ≈ 0.1 V) [34]. The distribution between the different hydrodimers was reported. R2 4

R1

3

a: R1 = R2 = R3 = H b: R1 = R2 = R3 = M

2 1 R3 90

Mixed hydrocoupling, in DMF, between simple 1,3-dienes, 90 (R1, R 2 , R 3 = H, Me), and alkenes activated by ester or keto groups has been attempted [154]. Yields of the MHC are usually poor (0–58%), the LHD of the activated alkene predominating. Again, yields of the MHC increase as the difference between the reduction potentials of the two components decreases, for example, no MHC was formed when ethyl 3,3-diphenylacrylate (Ph 2C = CHCO2Et; −1.70 V) or 40b (−1.89 V) was co-electrolyzed with butadiene, 90a (−2.66 V), or if cyclopentadiene (< −2.8 V) was co-electrolyzed with ethyl 3-methylcrotonate, 16b (Me2C=CHCO2Et; −2.46 V). The only combination giving reasonable yield (58%) of the MHC was co-electrolysis of 90a with 16b (−2.46 V) [154]. This indicates that electron transfer in solution is required for mixed coupling.

B.

INTRAMOLECULAR MIXED HYDROCOUPLINGS

Intramolecular mixed hydrocoupling was achieved for the system displayed in Equation 17.13 (dimethyl malonate added as a proton donor). Two stereoisomeric tricyclic products (1:1) were obtained (90% with EWG = COOMe and 23% with EWG = CN). For EWG = CN, higher yields and higher stereoselectivity were obtained using as supporting electrolyte LiClO4 (77% yield with 3:1 ratio of isomers) or Mg(ClO4)2 (62%, 11.4:1). The stereoselectivity is probably due to chelation of the metal cation with the carbonyl oxygen of the lactone and one of the cyano nitrogens, which favors the two functions being on the same face of the new ring [155]. Reduction of the monoester and of the mononitrile under identical conditions gives hydrogenation of the double bond of the lactone unit and no cyclization. This indicates that the α,β-unsaturated lactone is more easily reduced than the α,β-unsaturated monoester/mononitrile. The α,β-unsaturated dinitrile and diester are the electrophores in the cases where cyclization was achieved. EWG EWG EWG

EWG

MeCN, CH2(COOMe)2 O

EWG

O

O

O

O MeCN, CH2(COOEt)2 O

CPE, Hg cathode

O

O O +

MeOOC MeOOC

O

© 2016 by Taylor & Francis Group, LLC

(17.13)

H

H

MeOOC

O

+

Bu4NBr, CPE, Hg cathode

O EWG = COOMe, CN

EWG

1 : 1 Total yield 35–41%

(17.14)

660

Organic Electrochemistry

Under similar conditions, reduction of an α,β-unsaturated lactone linked to a linear α,βunsaturated ester gave modest yields of two isomeric spirolactones (Equation 17.14). CV of simple monofunctional compounds indicates that in this case the α,β-unsaturated ester function is the electrophore [156]. The regioselective coupling (β- to the ester group) and the presence of a proton donor (diethyl malonate) indicates that the probable reaction pathway is α-protonation of the radical anion, further reduction and Michael addition of the resulting β-anion to the lactone. The analogous α,β,γ,δ-unsaturated ester, on the other hand, did not undergo cyclization under these conditions [156], whereas the methyl ester of abscisic acid (Equation 17.15), which also contains an α,β-γ,δ-unsaturated ester function, undergoes cyclization upon reduction in MeCN (containing AcOH) [157]. OH

OH

O

MeCN, AcOH COOMe

CPE, Hg cathode

51% O

(17.15)

COOMe Mixture of stereoisomers

Cyclic voltammetry of monofunctional model compounds indicated that the unsaturated ester functionality is the one most easily reduced in the case of Equation 17.15. In the absence of the OH group, coupling was not observed [157] since enolization of the enone to the enol with conjugation of the four double bonds with the ester group takes place.

IV. REDUCTIVE COUPLINg OF ALKENES WITH OTHER TyPES OF SUbSTRATES A.

INTERMOLECULAR REDUCTIVE COUPLING OF ALKENES WITH OTHER TYPES OF SUBSTRATES

1. Coupling with CO2 Radical anions derived from α,β-diactivated alkenes such as dialkyl maleates, 80, react with CO2, and at high CO2 concentrations this process effectively competes with LHD formation [158,159]. Experimentally CV of 80b still shows a one-electron reduction peak in the presence of CO2 but higher scan rates are required to detect the radical anion on the reverse scan than in the absence of CO2. At the same time, a new reduction peak appears approximately 200 mV cathodic of the original reduction peak [160]. Preparative scale co-electrolysis of 80b and CO2 at the potential of the first reduction peak gave the dicarboxylated dimer, 91, whereas electrolysis at the potential of the “new” peak gave the dicarboxylated monomer, 92, in a 2 F process [160–162]. This can be interpreted according to the mechanism in Scheme 17.14, that is, the initially generated 80b−• adds rapidly to CO2 forming a distonic radical anion that is not reduced at the potential of the initial reduction. The distonic radical anion then dimerizes leading to the dianionic dicarboxylated dimer (which can be isolated after alkylation). Electron transfer between 53b−• and CO2 is too slow to be important because of the large difference (>0.4 V) between the reductions potentials of 53b and CO2. The mechanism in Scheme 17.14 seems to be general for α,β-diactivated alkenes in their reaction with CO2 [158]. The kinetics of the reaction between the radical anions of fumarates, 79a–c, and maleates, 80a–c, with CO2 in DMF (Bu4NI) has been studied by RRDE [159]. Reaction between the radical anion and CO2 prior to dimerization was confirmed for 80a by preparative and electroanalytical experiments at different substrate concentrations [162]. The radical anions 80−• reacted with CO2 20–50 times faster than did the analogous species 79−• for which the pseudo-first-order rate constants were determined to be in the range 0.35–1.5 s−1 [159]. Reduction of ethyl cinnamate in the presence of CO2 leads to the formation [168] of mono- and dicarboxyated adducts, 93–95. The overall yield and product distribution was found to be dependent upon the choice of electrode, potential, concentration of starting material, and temperature. The highest efficiency was achieved using a Ni cathode and a sacrificial Mg anode at a temperature

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661

Reductive Coupling – EWG EWG

–O2C 2

–O

+ CO2

EWG

2C

EWG

–O C 2

EWG

EWG EWG

EWG

EWG

CO2–

EWG 91

–O2C

EWG

EWG

e– 2 Peak

–O

2C

EWG –

EWG

CO2

–O

EWG

2C

CO2–

EWG 92

SCHEME 17.14 Reductive carboxylation with coupling.

of −10°C and a 0.1 M concentration of ethyl cinnamate. In contrast, a poor overall yield (30%) was obtained using a graphite electrode

Ph

CO2Et +CO2

–1.7 V, Ni cathode, Mg anode, CH3CN, 0.1 M Et4NBF4, 0.1 M substrate, –10°C (78%; 73:10:17 = ratio of 93:94:95)

Ph

– CO2 CO2Et + + CO2Et – Ph CO2 93 94

– CO2 CO2Et – CO2

Ph 95

Alkyl cinnamates, 15, are more easily reduced than CO2 by ≈0.4 V. Reduction of 15 in DMF (containing water) changes from a 1 F to a 2 F process in the presence of CO2. The radical anions react much faster than in the absence of CO2, and much faster than can be accounted for by electron transfer to CO2 [158]. On a preparative scale, the isolated products are mainly β-carboxylated monomer (see 94), together with the hydrogenation product. From methyl cinnamate, 15b, at low water concentrations, small amounts of the dicarboxylated monomer, 95, are also formed (Table 17.14), whereas with increasing water concentration, the amount of hydrogenation product increases at the expense of the β-carboxylated monomer [158]. This indicates that the H2O + CO2 equilibrium increases the acidity of the medium sufficiently for protonation of the radical anion to compete with both dimerization and nucleophilic addition to CO2. β-Carboxylation (rather than the expected α-) can be rationalized by assuming that the neutral radical arising from α-protonation of the radical anion is more easily reduced than the substrate [71], and the resulting β-carbanion adds to CO2 in competition with further protonation (Scheme 17.15), R = Ar. Like the reduction of ethyl cinnamate, the reduction of cinnamonitrile in the presence of CO2 leads to the formation of mono- and dicarboxyated adducts, 96–98. Once again, the overall yield and product distribution was found to be dependent upon the choice of electrode, potential, concentration of starting material, and temperature. The optimal combined yield is achieved [169] using a Ni cathode, Mg anode, a potential of −1.75 V (vs. Ag/AgI), and substrate concentrations in the range of 50–200 mM. In keeping with the fact that the solubility of CO2 increases as temperature decreases, the overall yield improves from ca. 75 to 85% as the reaction temperature is decreased from 15 to 0°C.

Ph

CN + CO2

–1.75 V, Ni cathode, Mg anode, CH3CN, 0.1 M Et4NBF4, 100 mM substrate, 0°C (85%; 38:8:54 = ratio of 96:97:98)

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Ph

– CO2 CN + CN – Ph CO2 96 97

– CO2

+ Ph

98

CN CO2

662

Organic Electrochemistry

TAbLE 17.14 Reductive Coupling of Alkenes and Alkynes with CO2a Substrateb

Conditions

E-PhCH=CHPh CH2=CHCN (11) (−2.14 V) 7

Products and yields

DMF, Bu4NI, CCE MeCN, Et4NOTs, CPEc CH3CH2CN, Bu4NBF4, 2.8 M H2O, CCE MeCN, Et4NOTs, CCEc DMF, Et4NClO4, H2O

92 (meso) 92% 91 41%d see 94 50%d

[29] [160] [160] [161] [158]

[161,162] [160] [161] [162] [162] [162] [173] [173] [173] [173] [173]

CH2=C(Me)CN (12c) (−2.31 V) E-PhCH=CHCOOMe (15b) (−1.78 V)e E-PhCH=CHCOOEt (15a) (−1.80)e

MeCN, Et4NOTs, CCEc DMF, Et4NClO4, H2O (0.056–2.78 M)

CH2=CHCOOMe (16a) (−2.10 V) 16a 16a CH2=C(Me)COOMe (−2.27 V) MeCH=CHCOOMe (−2.48 V) MeOCH=CHCOOMe (−2.67 V) Z-MeOOCCH=CHCOOMe (80b) (−1.53 V) 80b CH2=CHCOMe (38a) (−1.91 V) 39a (−1.91 V) [MeOOCCH=CH]2(CH2)2 [MeOOCCH=CH]2(CH2)3 [MeOOCCH=CH]2(CH2)4 H-Cα≡Cβ-Pr Pr-C≡C-Pr H-Cα≡Cβ-Ph Me-Cα≡Cβ-Ph Ph-C≡C-Ph

MeCN, Et4NOTs, CPEc MeCN, Et4NOTs, CCEc MeCN, Et4NOTs, CCEh MeCN, Et4NOTs, CCEc MeCN, Et4NOTs, CCEc MeCN, Et4NOTs, CCEc MeCN, Et4NOTs, CPE (new peak)c

see 91 41%d see 94, 2 F of 11, 91. No yields given see 91 28%d see 94 25–70%, 2 F of 15b 20–65%, 91 0 at electrode surface. Product: lower yield in MeCN/H2O Product: lower yield in MeCN/H2O Product: only 7% yield in MeCN Undiv. cell, Mg-anode, Ew = −1.2 V. Carboxylation at α- or β-carbon as indicated; see Equations 17.19 and 17.20. MeOOC

COOMe COOMe COOMe

MeOOC

COOMe COOMe

MeOOC

COOMe COOMe

COOMe COOMe

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663

Reductive Coupling Proton donor present (H2O+CO2) – EWG EWG + H+ R R – EWG EWG – EWG + + R R R R EWG –O C EWG – 2 + CO2 R R 95 EWG EWG + – + H R R

Proton donor absent (–H2O), high substrate conc. – EWG R EWG 2 – – R R EWG R EWG R EWG – 2 CO2 O C – 2 – CO2– EWG R EWG R 99

EWG

Proton donor absent (–H2O), R = Ar EWG – EWG + CO2 R CO–2 – R EWG EWG – EWG + + R CO–2 R R CO–2 R – EWG EWG O2C – + CO2 R CO–2 R CO–2

EWG

92 Proton donor absent (–H2O), R = H, or R = alkyl – EWG EWG + CO2 + CO2– R R EWG EWG – + CO– 2 – – O2C R – EWG R EWG O2C – + CO – 2 R OC CO– 2

2

92

SCHEME 17.15 The fate of radical anions generated from activated alkenes: protic, aprotic and CO2 conditions for follow up reactions.

Much smaller potential differences ( 6 ≫ 4 ≈ 7. Formation of five-membered rings is inhibited in the presence of a proton donor (dimethyl malonate) and cyclization to four- six- and seven-membered rings is completely prevented. However, these conditions have little effect on the formation of the three-membered rings [228]. The authors interpret the results of their preparative studies as a competition between protonation and intramolecular SN2 reaction on the radical anion stage. However, it seems more likely that the competition is between protonation and cyclization of the β-carbanion formed en route to hydrogenation. Initial, electrochemical reduction of β-dicarbonyl enol phosphates linked to an olefinic chain gives cleavage of the phosphate with formation of a vinylic radical [229–231]. Reduction in DMF takes place in the potential range −2.0 to −2.3 V and may lead to bicyclic products as illustrated in Scheme 17.23. The vinylic radical participates in a tandem cyclization either through a radical mechanism or through an anion mechanism (Scheme 17.23). The radical pathway seems the more likely, with reduction of the final bicyclic radical being facilitated by the electron-withdrawing group. The stereochemistry at this center is determined by preferred protonation from the least hindered side. Ketones and esters can both function as the activating group. Generally formation of six-membered rings is less favored than that of five-membered rings, and if R3 ≠ H the yields are also lowered [229]. More examples are given in [230,231]. O (EtO)2PO

R2

R1

DMF, Bu4NClO4

R2

R1

EtOOC

COOEt

+e–, –(EtO)2PO–2

COOEt

e–

R1 R3

R3

R3

R2

COOEt R2 –

R1

R3

EtOOC EtOOC –

e–

R1 R3

EtOOC

R1

R2

R2

R3

H+

H EtOOC

R1

R2

– R3

SCHEME 17.23 Tandem cyclization—vinylic radical or anion intermediate?

© 2016 by Taylor & Francis Group, LLC

R1 50–60%

R2 R3

681

Reductive Coupling

5. Coupling of Activated Alkenes with Organic Halides Formation of four- or five-membered rings follows reduction of dialkyl bromoalkylidenemalonates by cyclization to the β-carbon of the diactivated double bond (Equation 17.45) [232]. Br

COOR

DMF, Bu4NBr

COOR

Ew = –1.85 V

n

n = 1, 2

COOR n

COOR

(17.45)

60–80%

A plausible mechanism may involve C-Br cleavage through inter- or intramolecular electron transfer from the initially formed radical anion to the σ*-orbital. The radical so formed may add intramolecularly to the activated alkene function. An alternative, SN2 reaction between the alkene radical anion and the bromide would be expected to occur from the α-carbon of the activated alkene, which bears the higher charge density. 2+

2+

2+

2+

2+

N

N

N

N Ni N N H –0.70 V Ni(CR)2+

N

N

N

HN

N

HN

Ni

Ni N

N

–0.95 V Ni(tmc)2+

H N

HN

NH

HN

Ni

Ni

–1.38 V Ni(tet a)2+

–1.16 V

NH NH

Ni(cyclam)2+

Similarly to the intermolecular couplings between alkenes and halides (Section IV.A.4), yields of intramolecular couplings may be improved by using the same type of transition metal catalysts. Examples of some commonly used Ni(II)-catalyst precursors are shown above. Here and in Section IV.B.6 only a few examples of their application will be given, more can be found in Chapters 24 and 36. Six-membered rings are formed in reductions, catalyzed by electrogenerated Co(I)- or Ni(I)complexes, of α,β-unsaturated esters linked through the β-position to a chain bearing a bromine atom. The complexes include vitamin B12a (Co) and Ni(II)(tmc)2+ [233], which, as mentioned in Section IV.A.4, are both known to cleave organic halides to alkyl radicals rather than carbanions. Coupling therefore involves radical attack on the activated carbon–carbon double bond, in this case with 6-exo-cyclization (Equation 17.46). The yields are in the range 40–85% being higher for monoactivated alkenes than for diactivated ones. For diactivated alkenes (R1 being a carbonyl oxygen), the yields are lower (≈20%), which may be accounted for by competing, direct reduction of the alkene. The products were mixtures of stereoisomers except in a single case where the configuration at the double bond was cis. Here the trans product was formed exclusively (85%) [233]. Br

COOR

R3 R2

R3 COOR

DMF, Et4NClO4 O

R1

Red., Ni(l) or Co(l) cathode

R2

O

1

R

(17.46) 20–85%

Vitamin B12a has been used as catalyst precursor for the reductive coupling of 114 to 115 [207,234–236]. The electrochemical reduction at −1.5 V of the Co(III)-alkyl complex formed by oxidative addition of the bromide to the Co(I)-species (formed at ≈ −0.9 V) probably leads to the anion that may add to the double bond in competition with protonation. In DMF, reduction at −1.54 V gave the cyclized product 115 but with no trans/cis selectivity (Scheme 17.24) [234]. In a microemulsion of CTAB/1-pentanol/tetradecane/water (17.5/35/12.5/35 w/w%) with the same Co(I) catalyst, the same reaction led to 115 with a trans/cis ratio of 14:1, independently of

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682

Organic Electrochemistry O

O Br

Microemulsion

n

n

Co(I) cathode, Ew = –0.9 V or –1.5 V 114

115 n = 1, 90%, trans/cis = 14:1

DMF, LiClO4, NH4Br, Co(l) cathode Ew = –1.54 V O n

n = 1, 95% trans/cis = 1:1 n = 2, 70%

115

SCHEME 17.24 Competition between protonation and Michael addition.

whether the reduction was carried out at −0.9 V or at −1.5 V [207] (Scheme 17.24). The stereoselective formation of the trans-1-decalone in the microemulsion was ascribed to equilibration of the two isomers via OH− catalyzed keto-enol tautomerization [237]. For 114 with n = 3, reduction at −0.9 V gave little cyclization (20%) in DMF as well as in microemulsions, indicating that the radical is not significantly undergoing 5-endo-cyclization [235]. At −1.5 V, where the anion is formed, the competition between Michael addition and protonation was different for reaction in DMF and in microemulsions, giving 19% cyclization in DMF and 62–70% in microemulsions, indicating that the anion is formed in a water-free part of the microemulsion. 6. Coupling of Unactivated Alkenes with Organic Halides Unactivated alkenes give intramolecular coupling with halide functions only in catalyzed reactions since the low potentials required for direct reduction result in formation of highly reactive carbanionic intermediates. Even the direct reduction of aryl halides, linked to an ortho-alkene side chain, normally give the aryl anion rather than the radical since the potential of the initial reduction is lower than that required for furhter reduction of the phenyl radical. Cyclized products may be obtained but normally in low yields due to competing protonation of the phenyl anion, as demonstrated by addition of D2O to the solvent [238]. Outersphere electron transfer reagents may afford reduction at potentials where either the phenyl radical is stable toward further reduction or the rate of reduction by bimolecular reaction with the catalyst is so slow that intramolecular coupling or hydrogen abstraction competes. This has been tested for o-(3-butenyl)bromobenzene where the results of direct reduction (Ew = −2.65 V) were compared with reductions using m-tolunitrile as an electron transfer mediator (Ew = −2.25 V) [238]. The mediated reduction afforded a higher ratio of cyclized to uncyclized product (8:1 vs. 1:1) and, in contrast to direct reduction, the mediated reduction was little affected by added proton donor. Radical cyclization to the fivemembered ring was estimated to have a rate constant of ≥107 s−1 [238]. The carbon–fluorine bond of an aryl radical anion cleaves more slowly than does the carbon–bromine bond, and consequently aryl radical formation takes place in solution following direct reduction. In this case, a ratio of cyclized to uncyclized product of 2.5:1 was found for direct reduction [239] even in the presence of 0.5 M H2O. Another example is the direct reduction of 116a,b, where the presence of the cyano group facilitates reduction and which gave hydroindole coupling products (30–35%) (Equation 17.47). The major side products were the dehalogenated starting materials. Intramolecular coupling of the analogous 116c containing no cyano group was only obtained by indirect electrolysis using either anthracene or stilbene as mediator [240].

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683

Reductive Coupling COMe N

R2 R1

Cl

MeCN, Et4NBF4 Hg cathode

COMe N

R2 R1

(17.47)

116 1

a–b: 30% c: 80%, stilbene as catalyst

2

a: R = CN, R = H b: R1=H, R2=CN c: R1 = R2 = H

A variety of Ni(II)-complexes with macrocyclic tetradentate nitrogen ligands (see the examples in Section IV.B.5) has been applied as catalyst precursors for coupling of unactivated alkene functions with halides. The reactivity of the active, reduced Ni(I)-form varies with the reduction potential. Selected Ni-complexes can therefore be used to catalyze intramolecular reductive couplings for different types of halogen compounds with varying carbon–halogen bond strengths. The Ni-complexes have been used to perform intramolecular cyclizations to five-membered rings with a range of substrate types containing unactivated double bonds linked to various halides, for example, alkyl [241–243], vinyl [244,245], and aryl halides [243,246]. Standard conditions are reduction at the potential of the Ni(II)-complex (0.1–0.2 eq.), in DMF with tetraalkylammonium salts as supporting electrolyte. The appropriate Ni(II)-complex can be chosen as the one most easily reduced for which catalysis takes place on a voltammetric time scale [242]. Since the Ni(I)-catalysis gives radicals, the initial products of the 5-exo-cyclizations are also radicals, which (in contrast to those obtained by addition to activated alkenes) are normally not reduced under the conditions of the electrolysis. The follow-up processes are therefore radical reactions, normally assumed to be H-atom abstraction from the solvent. This has been demonstrated for 117a (Equation 17.48), where the yield of cyclized product was shown to be higher in DMF than in MeCN, DMF being a better H-atom donor than MeCN [242]. Cyclization from 117a in MeCN could be improved by addition of Ph 2PH as an external H-atom donor, otherwise the main product was the halogenated cyclized compound formed by halogen atom abstraction from another molecule of starting material. However, addition of Ph2PH also increased formation of acyclic products due to competing H-atom abstraction of the initially formed radical [242]. Using 117b rather than 117a and carrying out the electrolysis in MeCN without Ph2PH it was possible to obtain the iodo compound as the only cyclic product due to the lower activation energy for iodine radical abstraction than for bromine radical abstraction [243].

Red., Ni(l) cathode

a: X = Br b: X = I 117a, DMF: 117a, DMF, 2 eq. Ph2PH 117a, MeCN 117a, MeCN, 2 eq. Ph2PH 117b, MeCN

+ N

O

(17.48)

+ N

O

Ts

Ts

N Ts

41%

5%

14%

59%

8%

—

8%

—

33%

58%

15%

—

—

14%

61%

O

–

117

O

–

Ts

Et4NClO4 –

–

N

X

H

X

Under conditions similar to those in Equation 17.51 (DMF as solvent, divided cell), reduction of 118 (X = I) gives 75% of the cyclized product [241]. The yield was improved (90%) using an undivided cell and a sacrificial Mg anode (Equation 17.49) [246]. The yields increased in the order Cl < Br < I.

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684

Organic Electrochemistry X

DMF, Bu4NBF4 60–90%

Ni(l) cathode, undiv. cell, Mg anode

O 118

O

(17.49)

X = Cl, Br, l

The carbon–halogen bond is stronger in vinyl halides than in alkyl halides, and Ni-complexes with slightly lower reduction potentials have to be used to catalyze the reductive cyclization of vinyl halides on to alkene functions. Using conditions similar to those for the alkyl and aryl halides 5-exo-cyclized products were obtained from 119 [241]. In a single case (methyl substitution at the position of 5-exo-attack), only the 6-endo-product was obtained (Equation 17.50) [241]. R2

R2

Br R1

R3

DMF, Et4NClO4

R1

R3

Red., Ni(l) cathode MeOOC COOMe MeOOC 119

COOMe

MeOOC

COOMe

(17.50)

32%

45–85% R1 = H

R1 = Me

R2, R3 = H, Me

R2, R3 = H, H

Tandem cyclization was not obtained for 119. However, this could be achieved for 120 in which the vinyl halide was further activated with an acetyl group (Equation 17.51), and the catalyst was used in a larger amount (0.3–0.5 eq.) [245]. R1

R1 X

COMe

DMF, Et4NClO4

COMe

Red., Ni(l) cathode

Y

Y

40–50%

(17.51)

120 R1 = H, Me, Ph X = Br, l Y = C(COOEt)2,N-Ts, N-allyl

In explanation of the difference, it was proposed that the subsequent formation of the cyclopropyl ring by 3-exo-cyclization is reversible. The electron-withdrawing group at the final radical center promotes its rapid reduction (by the “excess” Ni(I)-complex) to a carbanion that is rapidly protonated [245]. The substrate, 121, has been cyclized using reduced vitamin B12a as a catalyst (0.01 eq.). The radical formed by cyclization can be trapped with an activated alkene (added in excess) to form a new radical that probably undergoes reduction and protonation to give a mixture of diastereoisomeric products [247] (Equation 17.52). EtO O Br

DMF, LiClO4 Red., Co(l) cathode

121

O

OEt

OEt CN OAc

O

(17.52) CN

63%

OAc

7. Coupling of Alkynes with Organic Halides Cyclization by direct reduction of alkynes linked to halides is usually ineffective. Phenyl substituted alkynes are more easily reduced than alkyl chlorides but more difficult to reduce than

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685

Reductive Coupling

alkyl bromides and iodides. Simple alkynes are more difficult to reduce than any of the halides. Consequently, direct reduction of 6-bromo-1-phenyl-1-hexyne at the potential of cleavage of the C–Br bond only affords 12% of benzylidenecyclopentane, whereas the corresponding chloride gives the saturated benzylcyclopentane in 45% yield [248] or, at low substrate concentrations (4⋅105

b

2⋅104

b

−1

[258]

[267]

[267]

[258, 267]

[280] [280]

b

[261] [261]

c

DMF, Bu4NBF4, 22°C, DCV

1.6⋅106 d

122a

DMF, Bu4NBF4, AcOH, 22°C, DCV

2.3⋅106

a

122a

DMF, Bu4NClO4, DCV DMF, Bu4NClO4, 25°C, CV, LSV

106

e

[258]

1.1⋅105

106

[274]

1,3-Benzenedicarbonitrile

≠ < 1.2 kcal ∆H dim ≠ mol−1, ∆Sdim < −26 −1 cal K mol−1 ≠ = −0.2 kcal ∆H dim ≠ mol−1, ∆Sdim = −30 −1 cal K mol−1

[266]

[258]

(Continued)

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688

Organic Electrochemistry

TAbLE 17.20 (Continued) Selected Rate and Equilibrium Constants for Dimerization of Radical Anions Derived from Substituted Aromatics Substrate

Conditions

kdim/M−1 s−1

Kdim/M−1

9-Fluoro-10cyanoanthracene (132)

CH3CH2CN, Bu4NPF6, CV

1.8.10

1,3,5-Trinitrobenzene (133) Acridine (138)

MeCN, Bu4NPF6, CV

1.8.105

DMF, Bu4NClO4, DCV NH3, CF3SO3K, 21°C, 8.1 bar

2.8.105

b

4.8.103

8.4⋅102

Quinoline (139)

a b c d e

Comments

References

∆H = 2.4 kcal ≠ mol , ∆Sdim = −26 cal K−1 mol−1

5

≠ dim −1

[271]

[273] [263] Ea(dim) = 12 kJ mol−1; A = 7.105 M−1 s−1

[282]

Dissociation reaction not taken into account. Not reversible on CV time scale. Radical anion stable on CV time scale but dimer dianion formed on preparative time scale. Irreversible due to protonation. Reversibility not considered.

For 122e, the effect on kdim and on Kdim induced by change of the solvent (Table 17.20) is modest ≠ and not easily interpreted [267]. The activation enthalpies for the dimerization step, ∆H dim , are small ( 5. The dimerization in acidic aqueous solution has been studied by a variety of electroanalytical methods [284]. Preparative scale reduction of 140a in aqueous solution at pH 0.6 gave the hydrodimer, 4,4′-tetrahydrobipyrimidine, in a clean 1 F process [285]. The product could subsequently be oxidized by KMnO4 to bipyrimidine. In DMF, the two substituted thio derivatives of pyrimidine, 140b and 140c, were found to dimerize upon reduction in the absence as well as in the presence of acid [286]. For 140b, however, the coulometric n-value was close to 0.5 in the absence of proton donors, indicating that the N-proton in 140b is sufficiently acidic to protonate either the radical anion or the dimer dianion, thereby leaving half of the substrate as the irreducible anionic form. In the presence of stoichiometric amounts of HClO4, the reduction takes place via the preprotonated substrate as for 140a [286]. The rate constant for dimerization of 140c−• (in the absence of acid) was found to be close to that for 140a−• [286].

B.

INTRAMOLECULAR COUPLINGS

The major reason for using transition metal catalysts in the intermolecular coupling of aryl halides is to minimize the competition from further reduction of the radical formed by carbon–halogen bond cleavage of the radical anion. Intramolecular reaction outruns further reduction when the σ-aryl radical is sterically placed to react with a second aryl group. The slower the cleavage reaction the further away from the electrode is the radical formed, which favors coupling in competition with further reduction. However, since coupling between phenyl radicals and aryl groups normally is slower than between phenyl radicals and alkenes, the same drawbacks as discussed in Section IV.B.6 pertain to these reactions. The scope of the reaction type has been explored in particular by Grimshaw and coworkers [287–297]. Since the rate constant for cleavage of the carbon–halogen bond decreases in the order I > Br > Cl > F, yields of coupling products are usually higher using chloro rather than bromo compounds. Most of the studies have been carried out by controlled potential electrolysis (Ew ≈ −2.1 V for the chloro compounds) in DMF using divided cells and Hg-cathodes [287–296].

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692

Organic Electrochemistry

Where an aryl halide is connected to the “receiving” aryl function by an amide linkage, as in Equation 17.56, two different types of cyclization occur, depending on where the initial radical attack takes place [287,289–291]. The relative yields of the two cyclization products depend on the substitution pattern in the aryl group being attacked [289]. Only the syn-conformation of the radical anion can give rise to coupling, and rotation about the amide bond is particularly slow when o-substituents are present. Where the aryl group under attack is o-disubstituted 95%. In the absence of added base, PhSO2− is probably the active but less efficient base leading only to 68% of 143 along with a hydrogenated derivative and the noncyclized 2 F product, 144 [301] (see Scheme 17.28). Although the electron transfer from 143− • to 144 is endothermic, the overall reaction is driven by the fast cleavage of 143− •. Intramolecular reductive coupling between aromatic rings (phenyl, naphthyl, anthryl) and nonconjugated keto groups have been achieved in 25–70% yield [302] (Equation 17.59). Lower yields were obtained with other combinations of solvents and electrolytes and by application of other cathode materials. O

H i-PrOH, Et4NOTs

Me

OH 70%

(17.59)

Sn cathode, div. cell, 4 F

However, only six-membered rings were formed successfully, and only nonaromatic ketones could be used [302]. The reaction is initiated by reduction of the keto group, and the reaction is highly stereoselective with a cis arrangement for the hydroxy group and the hydrogen at the coupling site. The neutral medium and observed effects of substituents in the aromatic ring indicates that the coupling involves an ion-paired radical anion and that the diastereoselectivity is controlled by coulombic interactions.

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694

Organic Electrochemistry

Compounds 145 are nonplanar and undergo 1 F reduction in MeCN around −2 V [303]. The radical anions derived from 145a–d undergo stereoselective cyclization according to Equation 17.60, whereas 145e–h give stable radical anions. The initial product, the cyclized radical anion, is airoxidized upon work-up giving the 0 F cyclized product, which is identical to the product obtained by photoexcitation [303]. The radical anions that fail to cyclize are those in which extended delocalization is not achieved by flattening of the ring systems or in which steric hindrance prevents co-planarity of the ring systems [303]. W

W Y

Z

e–

– Y

Z

X

a: W = X = O, Y = Z = Me b: W = X = O, Y = H, Z = Me c: W = S, X = O, Y = Z = Me d: W = S, X = NBz, Y = Z = Me

C.

a–d

X O

O

W X

O

O O 145

Y Z

e–h

– O2 –O2–

Y Z

O

W X

O

O

(17.60)

e : W = X = O, Y, Z = adamantylidene f : W = S, X = NPh, Y = Z = Me g: W = S, X = NC6H4NO2, Y = Z = Me h: W = S, X = NCH2COH, Y = Z = Me

COUPLING BETWEEN AROMATIC RINGS AND OTHER REAGENTS OCOCH3 DMF, Bu4Nl, Ac2O Hg cathode

66–75%

(17.61)

Reductive coupling, in aprotic solvents, of aromatic hydrocarbons with CO2 has long been known. Reduction of naphthalene in DMF in the presence of CO2 leads to 1,4-dicarboxy-1,4-dihydronaphthalene (≈50%) [304]. Under similar conditions, phenanthrene gives trans-9,10-dicarboxy-9, 10-dihydrophenanthrene (≈30%) [304]. Carboxylation of aromatic halides can be achieved in aprotic solvents either by using transition metal catalysts or by direct reduction using sacrificial anodes in undivided cells (see Chapters 24, 25, and 36). Reduction of anthracene in the presence of an excess of acetic anhydride led to formation of the enol acetate of 9-acetyl-9,10-dihydroanthracene (Equation 17.61) [305]. In each case, the hydrocarbon is the species reduced, but whether or not the ensuing radical anion is involved in electron transfer to CO2 or Ac2O prior to coupling was not investigated. Intermolecular coupling between aromatic rings and nonaromatic organic halides have been observed in a number of cases [306] and refs. therein. The mechanism of these reactions is well understood. The aromatic compound is reduced and electron transfer to the halide takes place in solution; the radical resulting from cleavage of the halide may either be reduced or undergo coupling with the aromatic radical anion depending on the relative reduction potentials of the radical and the aromatic system. In cases where coupling takes place, the anion formed in the coupling process may either be protonated by the medium or undergo SN2 reaction with another molecule of halide. An example of the first type of coupling reaction is the t-butylation of pyrene [307]. Upon reduction of pyrene in DMF in the presence of t-BuCl, 1-t-butylpyrene is formed as the major product (52%) after coupling, protonation and oxidation of the dihydro compound [307]. An example of the second type of coupling is the formation of 9,10-ethano-9,10-dihydroanthracene (57%) by reduction of anthracene in DMF in the presence of a moderate excess of 1,2-dichloroethane [308]. The reaction type is further discussed in Chapters 25 and 36. Intermolecular coupling of pyridines with ketones under acidic conditions has received attention, but like the intramolecular reductive coupling of aromatic rings linked to aliphatic ketones (Equation 17.59), the reaction is essentially a ketone reduction and the reader is referred to Chapter 31.

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D. REDUCTIVE DIMERIzATION OF POSITIVELY CHARGED AROMATIC SYSTEMS One-electron reduction of positively charged systems leads to neutral radicals, and independently of the structure the most common follow-up reaction is fast dimerization of two neutral radicals. In contrast to neutral radicals formed by cleavage, those formed by reduction of positively charged aromatic systems are not reduced at the potential where they are formed, and formation of 2 F products therefore only competes via H-atom abstraction. For N-heterocyclic systems in aqueous acidic media, protonation of the radical and further reduction leading to the dihydro-monomeric system may compete with dimerization. 1. N-Alkylpyridinium and Related Systems The electrochemical reduction of pyridinium cations and other positively charged N-heteroaromatic systems has received considerable attention as models for nicotinamide adenine dinucleotide (NAD+) and nicotinamide adenine dinucleotide phosphate (NADP+). The major reaction pathway in all cases consists of reduction to the neutral radical followed by coupling, primarily or exclusively in the 4-position with respect to the ring-nitrogen. Like the dimeric dianions formed by dimerization of, for example, 122 −•, the neutral dimeric products obtained from positively charged N-heteroaromatic systems are in many cases difficult to isolate since they are easily reoxidized back to the substrate cations. Reduction of the simple N-methylpyridinium ion, 146, is believed initially to give the expected N,N′-dimethyltetrahydro-4,4′-bipyridine but the end-product (in the absence of oxygen) is the N,N′dimethylbipyridine radical cation formed by a formal loss of two hydride ions and one-electron reduction of N,N′-dimethylbipyridinium [309,310]. The isolated product, N,N′-dimethylbipyridinium dication, results from air oxidation of the radical cation (Equation 17.62) [310]. kdim

2+

H N Me –

Me N –

N Me

N Me

–2 H–

H

–

–

N Me

Me N –

+

–

e–

(17.62)

e–

146

+

Me N

N Me –

–

CONH2 + N CH2Ph 147a

CONH2 + N CH2Ph 147b

CN + N CH2Ph 147c

+ N Me 148

N + 149

The reaction pathway in Equation 17.62 dominates for pyridinum ions unsubstituted in the 4-position, whereas 4-alkylpyridinium ions undergo reductive coupling to the 4,4′-tetrahydrobipyridine derivatives [311]. Values of kdim have been determined for several pyridinyl radicals in MeCN [311] (see Table 17.21). Despite the complication of air oxidation, the dimers have been isolated in 40–70% yield. Reductive dimerization of the NAD+-analogs 147a–c, in MeCN, and the reoxidation of the dimers have been studied in detail by CV and LSV [312]. This allowed the estimation of the Eo′-values and the dimerization rate constants (Table 17.21). Also in aqueous medium, 147a undergoes fast reductive dimerization, and the 4,4′-tetrahydrodimer was isolated as a mixture of two diastereoisomers [313]. The kinetics of the reductive dimerization of NAD+ in aqueous buffer (pH 9.1) has been studied by DPSC, CV, and LSV (see Table 17.21) [314]. Reduction of NADP+ on a preparative scale in 0.1 M NH4Cl/NH3 buffer (pH 9.3) leads to a mixture of dimers formed by coupling of the pyridine

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Organic Electrochemistry

TAbLE 17.21 Rate Constants for Reductive Coupling of N-Substituted Pyridinium, Quinolinium, and Acridinium Compounds Substrate

Conditions

N-Methylpyridinium (146) 1,2,4,6-Tetramethylpyridinium 2-ethyl-1,4,6-Trimethylpyridinium 4-ethyl-1,2,6-Trimethylpyridinium 2,6-diethyl-1,4-Dimethylpyridinium 2,4,6-Triethyl-1-methylpyridinium 1-Ethyl-2,4,6-trimethylpyridinium 1,4-Diethyl-2,6-dimethylpyridinium NAD+ 147a 147b 147c 148 148 a

kdim/M−1 s−1

H2O, KCl, pH 5–11, 25°C, polarogr. MeCN, Et4NClO4, CV MeCN, Et4NClO4, CV MeCN, Et4NClO4, CV MeCN, Et4NClO4, CV MeCN, Et4NClO4, CV MeCN, Et4NClO4, CV MeCN, Et4NClO4, CV H2O, Et4NCl, pH 9.1, 25°C, CV, LSV, DPSC MeCN, Et4NBF4, 20°C, CV, LSV MeCN, Et4NBF4, 20°C, CV, LSV MeCN, Et4NBF4, 20°C, CV, LSV MeCN, Et4NBF4, 20°C, CV, LSV MeCN, Et4NBF4, 20°C, DPSC

>10 2.5⋅106 2.5⋅106 1.8⋅106 2.6⋅106 1.5⋅106 2.0⋅106 8.0⋅105 3⋅107

Eo′/V

References

−1.372

7

−1.155

[317] [311] [311] [311] [311] [311] [311] [311] [314]

−1.105 −0.720 −0.520 −0.465 −0.460

[312] [312] [312] [312] [316]

a

7.9⋅108 2.0⋅108 5.0⋅108 2.5⋅107 (3 ± 1)⋅107

V vs. NHE. Strong adsorption of the dimer was found for C° > 2⋅10–4 M.

ring systems. The three diasteromeric 4,4′-dimers were the major products (≈70%) formed along with minor amounts of the 4,6′-dimers [315]. Reduction of N-methylacridinium ion, 148, in MeCN gives the expected dimer, 10,10′-dimethyl-9,9′-biacridine [312]. The value of kdim for 148 was found [312,316] to be about an order of magnitude smaller than that for the radical derived from 147 (see Table 17.21), probably due to the greater delocalization in the radical derived from 148. Reduction of acridizinium ion, 149, and substituted acridizinium ions in MeCN or DMF gives a dimer (≈80%) [318]. Although not confirmed experimentally, the most likely positions for coupling are indicated. The results of LSV measurements were in agreement with rate-determining dimerization of neutral radicals, and for 149 a lower limit for the rate constant of 107 M−1 s−1 was obtained by CV measurements [318]. The dimer could be quantitatively reoxidized to the substrate cations either electrochemically (at a potential ≈0.5 V anodic relative to the initial reduction peak) or by action of oxygen [318]. Rn + N

10% aq. H2SO4 O

n n = 1–3

Hg cathode

Rn

Rn H N n

OH +

H N n

Total yield OH

n = 1, 2; 40–62% n = 3, 15%

(17.63)

Inter- and intramolecular cross-couplings involving pyridinium rings are known. The crosscoupling of pyridine with acetone in acidic medium mentioned earlier is analogous to the intramolecular reductive cross-coupling of 1-(oxoalkyl)pyridinium ions (Equation 17.63). In contrast to the intramolecular coupling between phenyl groups and keto groups (Equation 17.59), which only gives rise to six-membered rings, good chemical yields are obtained of both five- and sixmembered rings in the case of pyridinium ions (whereas formation of seven-membered rings is inefficient) [319]. The reaction is 4 F overall since the pyridine ring also undergoes partial hydrogenation to form almost equal amounts of two tetrahydro isomers (Equation 17.63). The current efficiency is low due to competing proton reduction. The diastereoselectivity of the

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Reductive Coupling

reaction is high and opposite of that found in Equation 17.59 in which coupling was expected to involve the radical anion. The acidic medium used in Equation 17.63 ensures that the coupling involves the neutral radical formed by ketone reduction. The diastereoselectivity in this case has been explained by hydrogen-bond interaction between the OH-group and the ring nitrogen in the transition state leading to the preferred diastereoisomer and an unfavorable steric interaction between the ring and the alkyl group in the transition state leading to the other diastereoisomer [319]. Intramolecular reductive coupling between the isoquinolinium system, 150, and an attached aromatic halide (Equation 17.64 [320]) gives products similar to those obtained by the intramolecular coupling of aromatic halides with other aryl groups (Section V.B).

N+

MeCN, Et4NBr

I

Hg cathode

R R

N R

R = H, 86% R = MeO, 74%

(17.64)

R

150

2. Other Positively Charged Systems The pyranyl radicals formed by reduction of pyrylium cations, 151, dimerize at the 4-position, and anodic oxidation of the dimers leads to regeneration of the pyrylium cations [321] and references therein. Where the 4-position is unsubstituted, the dimerization process is very fast and irreversible, and a rate constant of 2.5⋅109 M−1 s−1 has been measured for the dimerization step in the reduction of 2,6-diphenylpyrylium cation, 151a, in MeCN by combination of LSV and fast CV using microelectrodes [322]. If a substituent other than H or Me is present in the 4-position, the dimerization process becomes reversible and in favor of the free radicals [321]. R4 R3 2

R

R5 + O 151

a: R2 = R6 = Ph, R3 = R4=R5 = H b: R2 = R6 = p-tolyl, R3 = R4 = R5 = H

R6

c: R2 = R6 = p-anisyl, R3 = R4 = R5 = H

In preparative scale, reduction of 151a in MeCN, bipyrilene, 153, has been observed as a side product in addition to the bi-4H-pyran, 152 [323]. Clean 1 F reduction was found (n = 0.8–1.0) in all cases, and at low substrate concentrations (5 mM) 152 was the only product. However, at substrate concentrations in the range 30–100 mM up to 48% of the product was 153 together with ≈15% 1,3-dibenzoylpropane. The conversion was not initiated by base or by the presence of oxygen, and the mechanistic explanation includes (Scheme 17.29) slow hydride transfer from the initially formed 152 to the substrate cation, followed by elimination of a proton and formation of 153. The 4H-pyran formed by hydride transfer reacts with water during work-up to give the dibenzoylpropane, and the apparent 1 F coulometry may be due to reduction of the liberated proton since the reduction was carried out using a Pt cathode [323]. Similar results were obtained in the reduction of 151b and 151c. The 1,2,3-triphenylcyclopropenyl cation, 154a, and cycloheptatrienyl cation, 155, give, like the heteroaromatic cations, free radicals by one-electron reduction, and these undergo almost quantitative dimerization in aprotic solvents [324,325]. Also, 154b undergoes reductive dimerization, the coupling position being exclusively at a phenyl substituted carbon atom [326]. After reduction of 154c in MeCN, only the rearranged dimer, 1,2,4,5-tetraphenylbenzene, was isolated (39%) together with N-acetyldiphenylcyclopropenylamine (24%) formed in a reaction with the solvent [324].

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698

Organic Electrochemistry Ar O Ar

Ar

Ar

H

+

O

+

Ar

H

O

+

O Ar Ar

Ar 152

Ar

Ar

O

O

Ar

O + H

H

H

Ar Ar

Ar

O

Ar Aq. workup

–H+

Ar CO (CH2)3 CO Ar

Ar 153

SCHEME 17.29 Cathodic reduction of pyrilium cations with fast irreversible follow up. Ph

Ph

Ph +

Ph

Ph

+

H

+

Ph

Ph 154a

+

Et

154b

154c

155

ACKNOWLEDgMENTS One of us (JHPU), who has impaired mobility, is especially grateful to Jessica Pancholi and Linda Malek (of SBCS) and Martin Beeson (QMUL Library) for help in accessing material.

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Organic Electrochemistry Hammerich, O.; Parker, V. D. Acta Chem. Scand. 1983, B37, 851–856. Koppang, M. D.; Woolsey, N. F.; Bartak, D. E. J. Am. Chem. Soc. 1985, 107, 4692–4700. Fawcett, W. R.; Lasia, A. Can. J. Chem. 1981, 59, 3256–3260. Gul’tyai, V. P.; Mendkovich, A. S.; Rubinskaya, T. Y.; Rusakov, A. I. Bull. Acad. Sci. USSR, Chem. Sci. 1990, 39, 1153–1155. Crooks, R. M.; Bard, A. J. J. Electroanal. Chem. 1988, 240, 253–279. O’Reilly, J. E.; Elving, P. J. J. Am. Chem. Soc. 1971, 93, 1871–1879. Navarro, I.; Rueda, M.; Ramirez, G.; Prieto, F. J. Electroanal. Chem. 1995, 384, 123–130. Tapolsky, G.; Robert, F.; Launay, J. P. New J. Chem. 1988, 12, 761–764. Battistuzzi, R.; Borsari, M.; Dallari, D.; Gavioli, G.; Tavagnacco, C.; Costa, G. J. Electroanal. Chem. 1994, 368, 227–234. Grimshaw, J.; Trocha-Grimshaw, J. Tetrahedron Lett. 1974, 993–996. Grimshaw, J.; Trocha-Grimshaw, J. Tetrahedron Lett. 1975, 2601–2602. Grimshaw, J.; Haslett, R. J.; Trocha-Grimshaw, J. J. Chem. Soc. Perkin Trans. 1 1977, 2448–2455. Grimshaw, J.; Mannus, D. J. Chem. Soc. Perkin Trans. 1 1977, 2456. Grimshaw, J.; Haslett, R. J. J. Chem. Soc. Perkin Trans. 1 1980, 657–660. Grimshaw, J.; Hamilton, R.; Trocha-Grimshaw, J. J. Chem. Soc. Perkin Trans. 1 1982, 229–234. Grimshaw, J.; Hewitt, S. A. Proc. Roy. Irish Acad. Sect. B 1983, 83, 93–101. Grimshaw, J.; Hewitt, S. A. J. Chem. Soc. Perkin Trans. 1 1990, 2995–2998. Donnelly, S.; Grimshaw, J.; Trocha-Grimshaw, J. J. Chem. Soc. Perkin Trans. 1 1993, 1557–1562. Donnelly, S.; Grimshaw, J.; Trocha-Grimshaw, J. J. Chem. Soc. Chem. Commun. 1994, 2171–2172. Donnelly, S.; Grimshaw, J.; Trocha-Grimshaw, J. Electrochim. Acta 1996, 41, 489–492. Vorozhtsov, G. N.; Dokunikhin, N. S.; Khmelnitskaya, E. Y.; Romanova, K. A. J. Org. Chem. USSR 1979, 15, 1744. Khmelnitskaya, E. Y.; Romanova, K. A.; Urman, Y. G.; Vorozhtsov, G. N. J. Gen. Chem. USSR 1980, 50, 2104–2108. Khmelnitskaya, E. Y.; Romanova, K. A.; Vorozhtsov, G. N. J. Gen. Chem. USSR 1981, 51, 1001–1006. Novi, M.; Dellerba, C.; Garbarino, G.; Petrillo, G. J. Chem. Soc. Chem. Commun. 1984, 1205–1207. Kise, N.; Suzumoto, T.; Shono, T. J. Org. Chem. 1994, 59, 1407–1413. Fox, M. A.; Hurst, J. R. J. Am. Chem. Soc. 1984, 106, 7626–7627. Wawzonek, S.; Wearring, D. J. Am. Chem. Soc. 1959, 81, 2067–2069. Lund, H. Acta Chem. Scand. 1977, B31, 424–438. Lund, H.; Daasbjerg, K.; Ochiallini, D.; Pedersen, S. U. Russ. J. Electrochem. 1995, 31, 939–947. Hansen, P. E.; Berg, A.; Lund, H. Acta Chem. Scand. 1976, B30, 267–270. Hobolth, E.; Lund, H. Acta Chem. Scand. 1977, B31, 395–398. Naarová, M.; Volke, J. Coll. Czeck. Chem. Commun. 1973, 38, 2670–2683. Yu, F. R.; Wang, Y. Y.; Wan, C. C. Electrochim. Acta 1985, 30, 1693–1701. Pragst, F.; Boche, E.; Koppel, H.; Walkhoff, E. J. Prakt. Chem. 1987, 329, 649–664. Anne, A.; Hapiot, P.; Moiroux, J.; Savéant, J.-M. J. Electroanal. Chem. 1992, 331, 959–970. Carelli, I.; Cardinali, M. E.; Moracci, F. M. J. Electroanal. Chem. 1980, 107, 391–404. Jensen, M. A.; Elving, P. J. Biochim. Biophys. Acta 1984, 764, 310–315. Ragg, E.; Scaglioni, L.; Mondelli, R.; Carelli, V.; Carelli, I.; Casini, A.; Finazziagro, A.; Liberatore, F.; Tortorella, S. Biochim. Biophys. Acta 1991, 1076, 37–48. Hapiot, P.; Moiroux, J.; Savéant, J.-M. J. Am. Chem. Soc. 1990, 112, 1337–1343. Gaudiello, J. G.; Larkin, D.; Rawn, J. D.; Sosnowski, J. J.; Bancroft, E. E.; Blount, H. N. J. Electroanal. Chem. 1982, 131, 203–214. Mitzner, R.; Bendig, J.; Ziebig, R.; Graichen, F.; Kreysig, D.; Pragst, F. J. Prakt. Chem. 1985, 327, 241–250. Gorny, R.; Schäfer, H. J.; Fröhlich, R. Angew. Chem. Int. Ed. 1995, 34, 2007–2009. Gottlieb, R.; Neumeyer, J. L. J. Am. Chem. Soc. 1976, 98, 7108–7109. Wintgens, V.; Pouliquen, J.; Kossanyi, J. New J. Chem. 1986, 10, 345–350. Amatore, C.; Jutand, A.; Pflüger, F. J. Electroanal. Chem. 1987, 218, 361–365. Pragst, F.; Seydewitz, U. J. Prakt. Chem. 1977, 319, 952–958. Shono, T.; Toda, T.; Oda, R. Tetrahedron Lett. 1970, 369–372. Breslow, R.; Drury, R. F. J. Am. Chem. Soc. 1974, 96, 4702–4703. Johnson, R. W.; Widlanski, T.; Breslow, R. Tetrahedron Lett. 1976, 4685–4686.

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18

Oxidative Coupling Hans J. Schäfer

CONTENTS I. Introduction ............................................................................................................................. 706 II. General Survey ....................................................................................................................... 706 III. Coupling via Radical Cations ................................................................................................. 707 A. Mechanism and Favorable Solvents ................................................................................ 707 B. Aromatic Compounds ..................................................................................................... 708 1. Hydrocarbons.......................................................................................................... 708 2. Phenols .................................................................................................................... 710 3. Aryl Ethers ............................................................................................................. 714 4. Aryl Amines ........................................................................................................... 714 C. Heterocycles .................................................................................................................... 716 D. Olefins ............................................................................................................................. 718 1. General ................................................................................................................... 718 2. Aryl Olefins ............................................................................................................ 718 3. Dienes ..................................................................................................................... 720 4. Enols and Enol Ethers............................................................................................. 720 5. Enamines ................................................................................................................ 721 6. Alkenes with Alkyl Substituents ............................................................................ 725 7. Alkynes................................................................................................................... 726 8. Alkanes ................................................................................................................... 726 E. Intramolecular Coupling ................................................................................................. 726 1. Aromatic Compounds............................................................................................. 726 2. Olefins ..................................................................................................................... 734 F. Anodically Induced Cycloadditions ................................................................................ 740 1. [2 + 2] Cycloadditions............................................................................................. 740 2. [4 + 2] Cycloadditions ............................................................................................. 742 IV. Coupling via Radicals ............................................................................................................. 747 A. General Comments ......................................................................................................... 747 B. Anodic Decarboxylation of Carboxylates as Radical Source ......................................... 748 1. Experimental Procedure ......................................................................................... 748 2. Homocoupling ........................................................................................................ 749 3. Heterocoupling ....................................................................................................... 752 4. Diastereoselective Coupling ................................................................................... 754 5. Addition of Radicals from Anodic Decarboxylation of Carboxylic Acids to Olefins ....756 6. Oxidation to Carbocations ...................................................................................... 761 C. Anions from CH-Acids as Radical Source ..................................................................... 761 1. Coupling ................................................................................................................. 761 2. Addition .................................................................................................................. 762 V. Outlook ................................................................................................................................... 765 References ...................................................................................................................................... 766

705

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Organic Electrochemistry

I. INTRODUCTION For the synthesis of organic molecules, two types of reaction are used: carbon–carbon bond–forming reactions and functional group interconversions. The former assembles larger molecules from smaller units; the latter introduces functional groups into a molecule or converts functional groups that are present in the molecule. The large number of carbon–carbon bond–forming reactions can be summarized in four groups: 1. 2. 3. 4.

Polar reactions between an activated acceptor and a donor or vice versa Radical coupling and addition reactions Pericyclic reactions, especially cycloadditions Transition metal catalyzed and organocatalyzed reactions

Electrochemistry can contribute to these synthetic transformations, especially to polar reactions, to radical reactions, and partially to cycloadditions. The reactive intermediates being involved, namely, radical ions, cations, and radicals, can be generated at the electrode simply, in large variety, cheap, and with reduced production of waste. There are several books and reviews that deal with the topic: organic electrosynthesis [1]. It should be added that organic electrosynthesis is in good accord with many rules of green chemistry [2]. The selectivity in organic electrosynthesis has been improved in the last decade by new developments in flow cells, microreactors, and reaction conditions (see Chapters 7 through 9) [1b,3]. This chapter deals with carbon–carbon bond–forming reactions at the anode and in laboratory scale.

II.

gENERAL SURVEy

Anodic coupling reactions may be arranged according to the reactive intermediates that are involved. These are radical cations that can undergo radical cation–radical cation coupling (Scheme 18.1, path a) or radical cation–substrate coupling with the starting compound (Scheme 18.1, path b). Radicals can be generated by deprotonation of C–H acids to form an anion, which is then oxidized. Furthermore, radicals can be formed by the one-electron oxidation of a compound R–H to a cation radical, which is subsequently deprotonated. The radicals can dimerize (Scheme 18.2, path a) or add to double bonds that subsequently lead to additive monomers 1 or additive dimers 2 (Scheme 18.2, path b). Furthermore, the radicals can be further oxidized to carbocations that can undergo electrophilic additions or substitutions with nucleophiles being in most cases nucleophiles that react at a heteroatom (Scheme 18.2, path c). These reactions can lead to carbon–carbon or carbon–heteroatom bond formation, and their course can be inter- or intramolecular. Furthermore, cycloadditions can be initiated by anodic generation of the dienophile or by inducing a chain reaction with a radical cation as initiator (see Chapter 15). a +

+ HR

+ RH

RH –e– RH

+

RH

R

b

+ HR

RH

SCHEME 18.1

–2H+

–e– RH

–2H+

Coupling reactions of electrogenerated radical cations.

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R

707

Oxidative Coupling a –e–

+

RH

–H+

R

R–R Y

R R

R

Y 1

R

RH

b –H+

R–

–e–

Y

c –e–

Y R

R+

HRNu, –H+

R–RNu

R Y 2

SCHEME 18.2

Coupling reactions of electrogenerated radicals.

The oxidation can occur in a direct electron transfer from the substrate to the anode or in an indirect oxidation via a chemical oxidant (mediator) being generated at the anode. The conversions are ordered in the first level according to reactive intermediates. However, one must keep in mind that this organization is not free from uncertainty since the mechanisms of anodic couplings are frequently just reasonable assumptions that are more or less supported by experiments. The mechanisms, shown in Schemes 18.1 and 18.2, illustrate additionally the principle of anodic umpolung [4]. In an electrochemical reaction, substrates of equal polarity can be coupled in a onepot reaction by combination of an electron transfer with a chemical reaction; nonelectrochemical reactions need at least two steps to achieve the coupling of substrates with equal polarity by applying chemical umpolung [5]. Furthermore by anodic umpolung synthetic building blocks can—in principle—be used in several reactivities, namely, by anodic oxidation of a nucleophilic building block to a radical, an electrophilic radical cation, or a carbocation.

III. COUPLINg VIA RADICAL CATIONS A.

MECHANISM AND FAVORABLE SOLVENTS

A review on the reactivity patterns of radical cations, which is supported by 562 references, describes reactions of a large variety of radical cations, which are prepared by chemical, electrochemical, and photochemical one-electron oxidation. The radical cations can undergo deprotonation, C–C or C–X bond cleavage, reaction with a nucleophile, cycloaddition, rearrangement, electron transfer, dimerization, radical reaction, or atom abstraction. Examples of these conversions are provided, and if available, thermochemical and kinetic data are given [6]. The stability of radical cations increases with increasing charge delocalization, blocking of reactive sites, and stabilization by electron-donating groups (hydroxy, alkoxy, and amino group) [7]. The complex reaction mechanisms of radical cations have been reviewed in detail [8]. Radical ions of conjugated systems can reversibly dimerize to form relatively stable σ-dimers as seen by cyclovoltammetric measurements [9]. Radical cations of 9-aryl- and 9,10-diarylanthracene derivatives with substituents in the 4-position of the aryl rings have been generated by photoionization in acetonitrile (AN). Their reactivity with n-butylamine and 1,4-diazabicyclo[2.2.2]octane and with anions as acetate, cyanide, bromide, and azide has been studied using nanosecond laser flash photolysis. The reactions proceed by electron transfer and/or nucleophilic addition. When electron transfer is thermodynamically feasible, this pathway dominates. For endothermic electron transfer, an inner sphere process can compete. The studies show that reactivity trends in the radical cation chemistry cannot be generalized as easily as those in carbocation chemistry [10]. The linear sweep voltammetry response to competitive radical ion–substrate coupling and radical ion dimerization mechanisms was determined by digital simulation. The simulations were carried out to mimic the conditions under which experimental studies had previously shown that the

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Organic Electrochemistry

radical ion–substrate coupling mechanism is the preferred reaction pathway. The configuration mixing (CM) model predicts an electronic reaction barrier for the radical ion–substrate coupling but not for the radical ion dimerization. The difference in standard free energy changes for the reactions is of the order of 7 kcal mol−1 or greater with the radical ion dimerization being energetically more favorable. In the discussed cases, the opposite result is observed experimentally: the radical ion–substrate coupling is the preferred pathway. The overall conclusion is that the CM model does not give reliable estimates of the reaction barriers for radical ion reactions [11]. In a review with 86 references, it is proposed that both polar and radical reactivity should be considered when discussing radical ion reactivity. In the past, the polar reactivity has dominated discussions. The following hypothesis is presented and supported by the literature: In the absence of severe steric effects, the reactivity of radical ions is dominated by the degree of coupling between charge and radical centers. Further work to test the validity of the hypothesis is proposed for many of the reaction types. It is concluded that the radical cation–carbenium ion comparison (for the reaction with acetate ion) would show similar reactivities. However, the radical cation–free radical comparison (for the reaction with dioxygen) would fail, since no reaction at all would be observed with the radical cation, while the free radical reacts rapidly [12]. Favorable reaction conditions for radical cations are provided by 1,1,1,3,3,3-hexafluoropropan2-ol (HFP). HFP is a solvent of low nucleophilicity, high hydrogen bonding donor strength, low hydrogen bonding acceptor strength, high polarity, and high ionizing power; this makes it an ideal solvent for radical cations in EPR spectroscopy, mechanistic studies, photochemistry, spin trapping, and synthesis [13a]. For the latter, see Section III.B.2. With HFP, the half-life of intermediate radical cations has been increased by a factor of 102 compared to trifluoroacetic acid (TFA) [13b].

B. AROMATIC COMPOUNDS 1. Hydrocarbons Alkyl-substituted aromatic hydrocarbons 3 can be coupled to diphenyls 6, probably by a radical cation– substrate or radical cation–radical cation coupling of the first formed radical cation 4 (Scheme 18.3, path a). Path b (Scheme 18.3, path b) leads to diphenylmethanes 7, probably by deprotonation of the R

3 –e– R R a 4,–2H+

R

b

R

6

4 –H+, –e– R

R

3,–H+ 7 5

SCHEME 18.3 Coupling of alkyl-substituted aromatic hydrocarbons 3 to diphenyl compounds 6 (path a) and diphenylmethanes 7 (path b).

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Oxidative Coupling

first formed radical cation 4 at the benzylic position; further oxidation of the benzyl radical leads to a benzyl cation 5 that afterward reacts in an electrophilic substitution with the starting compound. In the radical cation, a high positive charge density on an unsubstituted aryl carbon atom favors the diphenyl 6; a high positive charge density at the alkyl-substituted carbon atom supports the formation of diphenylmethane 7 [14]. To favor the coupling reaction, competing side reactions with nucleophiles must be suppressed by using an electrolyte of low nucleophilicity. A good choice is dichloromethane (DCM), tetrabutylammonium tetrafluoroborate (TBABF4), and small amounts of a strong acid to suppress a cathodic cleavage of DCM and subsequent chlorination, when electrolyzed in an undivided cell. Another solvent could be HFP that has been successfully used in phenol coupling (see Sections III.B.2 and III.B.3; there, HFP is named hexafluoroisopropanol [HFIP]). Applications are illustrated with some examples in Scheme 18.4. p-Xylene affords a diphenylmethane 8 in 22% yield at a carbon electrode in DCM, TFA, and TBABF4 [15]; 1,3,5-trimethylbenzene is dimerized exceptionally well in 71% yield at a platinum electrode in AN and TBABF4 to the diphenyl 9 [14]; and 1,2,4,5-tetramethylbenzene couples to the diphenylmethane 10 in 85% yield in DCM and TBABF4 [14]. Simpler conditions, namely, electrolysis in acetic acid, toluene, water and sodium benzenesulfonate in an undivided cell equipped with graphite felt electrodes, have been reported for the coupling of 2-methylnaphthalene to 2,2′-dimethyl-1,1′-binaphthyl and its isomer [16]. Cross-couplings are reported for naphthalene and pentamethylbenzene forming in 64% yield the diphenyl 11 (Pt, Bu4NBF4, AN, AcOH) [17] and for anthracene and anisole yielding 70% of anthracene 12 (Pt, Bu4NBF4, DCM, TFA, trifluoroacetic acid anhydride [TFAn]) [18]. Further examples can be found in References 17 and 19. The cyanation of aromatic hydrocarbons, which is also a C–C coupling reaction, is achieved with anthracene in AN, Et4NCN to yield 54% of 9,10-dicyanoanthracene [20]. The procedure is simplified by electrolysis in an emulsion system (water, DCM, NaCN, TBAHSO4) [21]. In this mode, 4-alkoxy-4-cyanobiphenyls, a class of liquid crystals, have been prepared. In an indirect anodic oxidation with Pd(OAc)2 and 2,2,6,6-tetramethylpiperidin-1-yl)oxyl (TEMPO), a stable radical as mediators arylboronic acids or arylboronates gave the corresponding biaryls in moderate to excellent yields [22]. Electrochemical oxidation of anisole, mesitylene, naphthalene, and anthracene in various ionic liquids (ILs) provided coupling products in an undivided cell. With 1,2-dimethoxybenzene (veratrole), a stable blue-colored doped polyveratrole was obtained [23].

2

8

9

10 OCH3

12 11

OCH3

SCHEME 18.4 Diphenylmethanes 8 and 10 and diphenyls 9, 11, and 12 in the anodic coupling of alkylsubstituted aromatic hydrocarbons.

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Organic Electrochemistry R1

R1

R1

R1

–H+, –e– 2

R3

R

R3 + R3

R2

OH 13

O 14

2

O

R1

+

+ R3

R2

R2

O 16 R3

O 15

R1

OH 17

R3 R2

–e–

–e–

R2

R2

R1 R1

R1

Nu

Nu– R3

O 18

+ R2

R3

R3

R2

O

O

Nu

19

Alkene [5 + 2] and [3 + 2] cycloadducts

Diene [4 + 2] cycloadducts

SCHEME 18.5 Anodic conversion of phenol 13 via phenoxy radicals 14, phenol radical cations 17, and phenoxonium cations 18. (Adapted from Yamamura, S. and Nishiyama, S., Synlett, 4, 533, 2002.)

2. Phenols Earlier literature on anodic coupling of phenols can be found in Reference 24. Furthermore, comprehensive, more recent reviews, which describe the anodic and chemical oxidation of phenols and phenol ethers, have appeared [25]. The oxidative conversion of phenols 13 can be described as shown in Scheme 18.5. From phenol 13, reactive intermediates as phenoxy radicals 14, phenol radical cations 17, and phenoxonium cations 18 are formed by electron transfer and deprotonation. The radical 14 can undergo C–C coupling to 15 and the p,p′- and o,p′-isomer of 15 and C–O coupling to 16 and the p,p′- and o,p′-isomer of 16. The cation 18 can react with alkenes to form [5 + 2] and [3 + 2] cycloadducts or with nucleophiles to form 19, which can produce [4 + 2] cycloadducts with dienes. If R3 = H, the dienone tautomerizes to the phenol. Radical cations 17 can undergo a C–C coupling or can be deprotonated to 14. Selected examples for the C–C coupling of phenols are shown in Scheme 18.6. The phenoxy radicals react by carbon–carbon and carbon–oxygen coupling to dimers that can be further oxidized, and in this way, product mixtures can arise. Blocking of the 2-, 4-, or 6-position makes the coupling reaction more selective. Phenols with an unsubstituted p-position usually form p,p′-coupling products as the major dimer (Scheme 18.6). From the corresponding phenols, the p,p′-dimers have been obtained: 20 (at a nickelhydroxide electrode and at 70°C) [26]; 21 in MeOH, DCM, LiClO4, at a Pt anode [27]; 22 (R = Me, Et, benzyl, i-Pr, t-Bu, CH2CH(CH3)CO2H) at a glassy carbon anode in MeCN (AN) and NaClO4 [28]; 23 at a graphite felt anode in MeONa, AN, and Et4NClO4 [29]; 24 at a Pt anode in MeOH, LiClO4, and NaOH [30]; and 25 at a Pt anode in AN, and NaClO4, Et4NOH [31]. If the 4-position is blocked, o,o′-coupling becomes the main reaction as found in 26–30 in HFIP (HFIP = HFP) and at a boron-doped-diamond anode (BDD anode) [32]. When the o- and p-position are blocked, an addition can occur as found in the cross-coupling between 2,6-di-t-butyl-cresol and anisole to form 31 [33]. In a template-directed anodic phenol-coupling reaction, eight different phenols were converted into the corresponding sodium (tetraphenoxy)borates. Their controlled-current electrolysis (CCE) in an undivided cell at a platinum electrode in AN afforded 2,2′-diphenols in 85–20% yield. This selective ortho-coupling reaction can be performed on a kilogram scale [34]. Using the BDD electrode (BDD anode) and hexafluoroisopropanol as additive, several phenols could be coupled to the o,o′-dimers in good to moderate yields (26–30, Scheme 18.6) [32,35].

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711

Oxidative Coupling O

OH

t-Bu

OH

t-Bu t-Bu

t-Bu MeO

CH=NR

MeO NMe HO

2 20 (90%)

2 21 (93%)

2 Me 23 (69%)

2 22 (41–100%)

CH3

CHO HO MeO

2 OH 24 (100%)

MeO

2 OH 25 (65%)

O O

2

2

26 (74%)

CH3 OH 27 (47%)

O CH3

CH3

Cl

2

Br

2

OH

OH 29 (30%)

28 (30%)

SCHEME 18.6

t-Bu

t-Bu

CH3

HO 2

OCH3

30 (41%)

31 (81%)

Selected examples for the anodic coupling of different phenols.

The two concepts of template-directed coupling and use of the BDD anode have been reviewed [36]. The procedure using the BDD anode was further developed to allow the application of a less expensive electrode (graphite) and a fluorinated solvent (TFA), both permitting the selective o,o′coupling in broad scope [37]. The cross-coupling of 4-methylguajacol with different arylethers affords at the BBD anode unsymmetrical biaryls in a one-step reaction without using activating auxiliaries. Depending on the reaction conditions, selectivities of unsymmetrical to symmetrical coupling product of 1.5:1 to >50:1 with yields of 16–47% are obtained [38]. In a very recently reported cross-coupling reaction, phenol (5 mmol) and arene (15 mmol) have been oxidized in N-methyl-N,N,N-triethylammonium methylsulfate, HFIP, and methanol in an undivided cell at a platinum anode and a nickel cathode. At a current density of 2.8 mA cm−2, the heterocoupling product was obtained in 67% yield and 69% current efficiency; the ratio of unsymmetrical to symmetrical product was >100:1 (Scheme 18.7). The large scope of the reaction is demonstrated with 13 different phenols and 8 different arenes, and partly selectivities >100:1; a mechanism is proposed [39]. Anodic oxidation of 2,4-dimethylphenol at platinum electrodes and Ba(OH)2·8H2O in methanol as electrolyte provides in an undivided cell a pentacyclic dehydrotetramer as single diastereoisomer in up to 24 g per run (52% yield) [40]. This dehydrotetramer was used to create a variety of polycyclic scaffolds that partially occur as core structures in natural products [41]. The 2,4-dimethylphenoxy group could be substituted by several amines; there primary amines provided the best yields [42]. OCH3

OCH3 BDD anode, –e– OH

Et3NMe+ –O3SOMe HFIP, MeOH

OH

+ CH3O

OCH3 OCH3 CH3O OCH3

SCHEME 18.7

OCH3

Highly selective cross-coupling of 4-methylguajacol with 1,2,4-trimethoxybenzene.

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Organic Electrochemistry

Electrochemical applications of the BDD electrode have been reviewed [43]. The electrochemical activation of the BDD electrode for electroanalytical investigations in organic media is described [44]. The electroanalytical oxidation of different phenols at the BDD anode and the glassy carbon electrode in aqueous and methanolic media and in microemulsions has been compared [45]. A recent review covers the applications of the BDD electrode in electroorganic synthesis. It reports cathodic carboxylations and reductions of oximes. At the anode, alkoxylation, fluorination, cyanation, C–C bond cleavage, and phenol coupling are described. Furthermore, informations on the stability of the BDD electrode and on electrolysis cells for BDD electrodes are communicated [46]. In the products shown in Scheme 18.5, by variation of the phenol structure and reaction conditions as oxidation potential, solvent—supporting electrolyte, electrode material, partially good selectivities toward single products can be achieved, and these have been used to synthesize complex natural products in few steps. An example in Scheme 18.8 shows that substituents can efficiently control the mode of coupling. The dibromo derivative of N-protected l-tyrosine is anodically oxidized in a current controlled electrolysis (CCE: 0.11 mA cm−2, LiClO4, MeOH) to afford 45% of the C–O coupling product that is deprotected quantitatively to iso-dityrosine [47]. With iodo instead of the bromo substituents under the same reaction conditions, 28% of the C–C coupling product was obtained that was deprotected in 72% yield to dityrosine. Scheme 18.9 shows an intramolecular electrophilic addition of a presumably anisole radical cation to a nucleophilic enol ether to form 51% of a spiro compound (CCE: 11.2 mA, LiClO4, 20% MeOH, DCM, 2,6-lutidine, reticulated vitreous carbon [RVC] anode) [48]. Similarly, 4-(2-alkenylaryl)phenols have been cyclized to the corresponding spiro compounds. There, the oxidation at the anode has been compared with the use of iodobenzene diacetate as oxidant. Yields up to 92% (PhI(OAc)2, MeOH)) and 80% (anode, AN, MeOH) have been obtained. With the exception of two cases, the same phenols, which gave good yields of spiro dienones in the anodic oxidation, gave also good yields with iodobenzene diacetate [49]. CO2Me 1. –e–, MeOH, LiClO4 2. Zn/HOAc CO2Me

NHZ Br

CO2Me CO2Me

O

Br OH

45%

NHZ

Br

CO2Me X

1. –e–, MeOH, LiClO4 2. Zn/HOAc

X OH

NHZ

X = Br or I I

2 OH 28%

SCHEME 18.8

C–O or C–C coupling controlled by the bromo or iodo substituent.

–e–, 51% O

MeO

MeO

OMe

MeO

SCHEME 18.9 Intramolecular reaction of a presumably anisole radical cation with a nucleophilic enol ether to a spiro compound.

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713

Oxidative Coupling OCH3 CPE, –e– AcOH, MeNO2 O

HO

OCH3 +

O

OCH3

O OCH3

OCH3

H

H O

O

O H

O

H O

SCHEME 18.10 Anodically initiated [3 + 2] cycloadditions between methyl hydroquinone and alkenes; CPE = controlled potential electrolysis.

Formal [3 + 2] cycloadditions were observed in the oxidation of methyl hydroquinone in the presence of styrene, dihydrofuran, and dihydropyran in 95%, 11%, and 33% yield, respectively (Scheme 18.10) [50]. In the anodic oxidation of 2-methyl-4,5-dimethoxyphenol and 1-(3,4-dimethoxyphenyl)-1propene, a [5 + 2] cycloadduct of the phenoxonium cation is obtained in 80% yield (Scheme 18.11) [51]. For additional anodic cycloadditions of phenols, see Section III.E.2. The anodic oxidation of 2,6-di-t-butylphenol was studied in Bmim ILs at platinum electrodes [52]. The electrolysis led to the diphenoquinone; however, better results were achieved with the electrolytes: methanol, DCM, and LiClO4. Phenol coupling can be achieved by Michael additions to o-quinones that are performed by in situ anodic oxidation of catechol in the presence of Michael donors (Scheme 18.12) [53]. For this O OCH3 OCH3

CCE, –e–

+ CH3

CH3

CH3

OCH3

Ac2O,HOAc Bu4NBF4

MeO OH

O

OMe MeO OMe

SCHEME 18.11 [5 + 2] cycloaddition in the anodic oxidation of 2-methyl-4,5-dimethoxyphenol and 1-(3,4-dimethoxyphenyl)-1-propene. OH OH H2O, NaOAc

OH + OH

SCHEME 18.12

O

O

–e–, 1.1 V 95%

O

Anodic cross-coupling of catechol with a Michael donor.

© 2016 by Taylor & Francis Group, LLC

OH

O

O

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Organic Electrochemistry

reaction, a wide scope has been demonstrated with differently substituted 1,2- dihydroxybenzenes and different Michael donors [54]. The anodic oxidation of catechols in the presence of α-oxo heterocyclic ketene N,N-acetals was investigated by using cyclic voltammetry and controlled-potential electrolysis. The results indicate that α-oxo heterocyclic ketene N,N-acetals can undergo a Michael addition at the anodically generated o-benzoquinones and selectively form arylated products in good yields. In addition calculations were performed to explain the exclusive formation of these products [55]. For further details on the coupling of phenols at the anode, see Chapter 26. 3. Aryl Ethers The anodic coupling of aryl ethers has been reviewed before [56]. Aryl ethers are more selectively coupled than phenols. The C–O coupling is excluded; the o-coupling and the oxidation to quinones become more difficult. A mixture of TFA and DCM proves to be a very suitable electrolyte [57]. Some selected examples are shown in Scheme 18.13. Anisol was dimerized in TFA, DCM (1:2), Bu4NBF4 at a Pt anode to diphenyl 32 (63%) [57], 1,2-dimethoxybenzene was converted under the same conditions to the dimer 33 (86%) [57]. 9-Methoxyanthracene afforded in AN, TFA, and Bu4NBF4 the dimer 34 (95%) [58], and 1,2-dimethoxybenzene led in DCM, TFA, and Bu4NBF4 to the cyclotrimer 35 (50%) [59]. 9-Anisyl-anthracene formed with anisole in DCM, AN, TFA, and Bu4NBF4 the cross-coupling product 36 (90%) [60]. From a catechol ketal in AN and Bu4NBF4, the cyclotrimer 37 (90%) was obtained in a 20 g scale; it was applied as platform structure for rigid receptors [61]. From 4-t-butyl-anisole in 0.1 M TBABF4 in AN, an o,o′-dimer (33%) and a trimer (20%) with an o,m′-coupling of 4-t-butyl-anisole to the dimer have been obtained [62]. 4. Aryl Amines The anodic coupling of aryl amines leads to three major products: benzidines, aminodiphenylamines, and azo compounds. The first intermediate in these oxidations is mostly the radical cation OCH3 OCH3

OCH3

OCH3

OCH3 OCH3

2

2

32 (63%)

33 (86%)

CH3O CH3O

2 34 (95%)

35 (50%)

OCH3 CH3O CO2Et

OCH3 O O O EtO2C O

37 (90%)

O O

OCH3

36 (90%) CO2Et

SCHEME 18.13

Intermolecular anodic coupling of aryl ethers to dimers and trimers.

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Oxidative Coupling

NR2

2

R2N

–2e–

a –2H+ +

2

+ o,p– + o, o-product R N

b –2H+

NR2

NR2

NHR

+ o -product

+ c –2H

R R N N

SCHEME 18.14 Anodic coupling of an aryl amine to benzidine, aminodiphenylmethane, and hydrazobenzene. CH3

OMe OMe

NR2

CH3 NH

N

2 R: C2H5, C4H9 38 (40–50%)

CH3

OMe 2

CH3 40 (90%)

39 (42%)

N H

N

N

OMe

t-Bu N N t-Bu 42 (25%)

O

N N N N

O

43 (80%)

41 (98%)

SCHEME 18.15

Anodic coupling of amines.

of the amine. This can couple in three different ways (Scheme 18.14): (a) In a C–C coupling reaction, benzidines are formed; (b) in a C–N coupling reaction, aminodiphenylamines are obtained; and (c) in an N–N coupling reaction, hydrazobenzenes are produced. These compounds are often easily further oxidized. Earlier work is reviewed by Adams [7]. Some amine oxidations are presented in Scheme 18.15. Diethyl- and dibutylaniline are coupled in AN and Et4NClO4 to benzidine 38 in moderate yield, while with dimethylaniline a complex product mixture is obtained [63]; under the same conditions, the benzidine 39 is formed from a substituted naphthylamine, instead of the desired cyclization product [64]. Electrolysis of di-p-tolylamine affords in AN, Et4NCN in 90% yield the hydrazine 40; with lutidine as base, however, a dihydrophenazine is formed as major product [65]; 1,2,3,4-tetrahydrocarbazole affords in AN, H2O, and LiClO4 nearly quantitatively the coupling product 41 by substitution at atom C1 of the carbazole [66]; the mechanism of this interesting reaction has been investigated using electroanalytical techniques [67]. t-Butylamine has been deprotonated with lithium cyclohexylamide, and the resulting amide was oxidized in tetrahydrofuran (THF) and LiClO4 to the azo compound 42 [68]. At the nickelhydroxide electrode, N-amino-morpholine was converted in 80% yield to the azo compound 43 [69]. The combination of electrochemistry with organocatalysis allows the substitution of a 4-hydroxyaniline in m-position; this selectivity cannot be achieved in a Friedel–Crafts reaction. Thereby, an intermediate enamine, formed from an aldehyde and a secondary amine, undergoes a Michael addition to an iminoquinone that results from an anodic oxidation of the 4-hydroxyaniline. The electrolysis is conducted current controlled in an undivided cell in AN, H2O, and

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Organic Electrochemistry NHTs

NHTs

Ph RCH — —O +

–e–

Ph + N H

OTMS

R OH

O OH

SCHEME 18.16 Regio- and enantioselective m-substitution of 4-hydroxyaniline by combining anodic oxidation with organocatalysis.

0.1M  NaClO4. With  a  chiral amine and five different aldehydes, the corresponding products are obtained in 69–87% yield and 81–96% enantioselectivity. Similar yields can be also obtained with iodosobenzene diacetate as oxidant (Scheme 18.16) [70]. DFT calculations have been used to determine theoretical values for the oxidation potentials for indole and carbazole derivatives. For an oxidative indole trimerization, the computed electron spin distributions of the involved radical cations support the proposed mechanism [71].

C.

HETEROCYCLES

Anodic syntheses of heterocyclic compounds dealing with the formation of C–N, C–O, and C–S bonds have been reviewed [72]. Cyclic amines as pyrrolidines, piperidines, tetrahydroquinolines, and tetrahydroisoquinolines can undergo anodic C–C coupling with the cyanide anion to form nitriles in α-position to the amino group [73]. Thereby, the nitrogen atom is oxidized to a radical cation that is deprotonated at the α-carbon atom, which is further oxidized to an amino methyl cation that subsequently reacts with the cyanide anion. For further anodic cyanations of heterocycles, see Reference 74. 1,3-Dithioles can be dimerized in AN and pyridine in moderate to low yields to tetrathiofulvalenes (Scheme 18.17) [75]. 2-Thiobarbituric acids can be coupled under ultrasonication in AN and DCM, 18-crown-6 in 10–27% yield to the corresponding dimers (Scheme 18.18) [76]. The electrochemical oxidation of 2,4-dimethyl-3-ethylpyrrole in AN was studied using cyclic voltammetry, constant current coulometry, preparative electrolysis, and ab initio calculations.

–e– AN, pyr

S

S

S

S

S

S 40%

SCHEME 18.17 Anodic dimerization of a 1,3-dithiole to a tetrathiafulvalene. S R2 H

R1

O R2

O N

N

–e–, K2CO3 AN, DCM, Pt

R2

S R1 = C12, C14, C16, R2 = CH3

N

R2

N

O

O R1

R1 O R2

O N

N S

SCHEME 18.18

Anodic dimerization of thiobarbituric acid.

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R2

717

Oxidative Coupling N

H

1. –e–(1.1 V); 2. +e–(0.0 V SCE)

OH MeO

AN, TFA, LiClO4, 60%

N

OAc H CH3 CO2Me CH3 MeO2C H AcO N

OMe

N H

OH H

SCHEME 18.19

OH N

MeO

N

OAc H CH3 CO2Me

Anodic dimerization of vindoline to 10,10′-bisvindoline.

The major product from preparative electrolysis was a trimer, for which a central 2H-pyrrole unit was proposed. Since 2H-pyrroles are stronger bases than the corresponding 1H-pyrroles, the trimer is effectively protected against further oxidation by protonation [77]. Pyrrole and thiophene derivatives can be polymerized in non nucleophilic solvents as AN, 1,2-propylene carbonate (PC), or DCM to conducting polymer films [78]. The anodic oxidation of catharanthine in the presence of vindoline was performed in AN and Et4NClO4 at a controlled potential to yield the alkaloid anhydrovinblastine. In this elegant synthesis, it is concluded from electroanalytical studies that the radical cation of catharanthine undergoes a fragmentation to a distonic radical cation, which after oxidation to a dication reacts in an electrophilic aromatic substitution with vindoline to afford the alkaloid [79]. Anodic oxidation of vindoline, at a Pt electrode and a controlled potential of E = 1.1 V versus SCE followed by reduction of the formed intermediate at E = 0.0 V versus SCE gave 10,10′-bisvindoline in 60% yield (Scheme 18.19) [80]. Cyclo[8]pyrrole was obtained efficiently, when 3,3′,4,4′-tetraethylbipyrrole was subjected to bulk electrolysis. The yields varied from close to 0% with Bu4NF as the electrolyte, to almost 70% when Bu4NSO4H was used for this purpose. These observations are consistent with the conclusion that the reaction is controlled by anion-related factors such as a specific templating effect (Scheme 18.20) [81].

–e– N H

N H

3 N H NH

N

SCHEME 18.20

2H+

N HN

HN

Anodic cyclotetramerization of 3,3′,4,4′-tetraethylbipyrrole to cyclo[8]pyrrole.

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Organic Electrochemistry

Carbazoles were successfully synthesized by oxidative cyclization of 2-aryl-N-acetylanilines using electrochemically generated hypervalent iodine as oxidant. The electron-withdrawing nitro group and the electron-donating methoxy group at the para-position of the acetamide group interfered with the cyclization. For the preparation of the oxidant, iodobenzene in trifluoroethanol (TFE), LiClO4 was electrolyzed at a glassy carbon beaker anode. After electrolysis, the substrate was added, and after stirring, 83% (97% based on conversion) of the carbazole was obtained. Electrolysis without the iodobenzene gave no carbazole [82].

D. OLEFINS 1. general Olefins with electron-donating substituents as the aryl, alkoxy, N-acylamino, thioalkyl, or vinyl group can be coupled in methanol to give 1,4-dimethoxy-dimers and/or dienes (Scheme 18.21). The first intermediate in this coupling reaction is a radical cation, which by a radical cation–substrate coupling and subsequent 1e-oxidation (path a) or by radical cation–radical cation coupling (path b) leads to a dimer dication. This undergoes a reaction with nucleophiles, in many cases a methanolysis and/or deprotonation. A further competing reaction is the reaction of the radical cation with a nucleophile from the electrolyte, which is methanol in many cases. This is followed by a further oxidation of the radical to a carbenium ion that reacts with a nucleophile. Thereby, an addition product of two nucleophiles to the olefin monomer is formed (Scheme 18.21, path c). 2. Aryl Olefins Styrene and indene derivatives (Scheme 18.21, Y = Ph) are dimerized to 1,4-diphenylbutanes or 1,4-diphenylbutadienes (Scheme 18.22) [83,84]. The product distribution depends in some cases on the anode potential and the supporting electrolyte. By substitution of sodium perchlorate for sodium camphorsulfonate or by increasing the anode potential, olefins are formed at the expense of methyl ethers. This could be due to an increased displacement of methanol from the anode surface by changing the more hydrophilic perchlorate anion to the more lipophilic camphorsulfonate anion. A similar trend is found in intramolecular cyclizations (see also Section III.E.2). Methanol depletion at higher anode potential might be due to the increased adsorption of anions with increasing potential. The surface concentrations of the reactants (solvent and styrene) govern the relative rates for the pathways a–c, which determine the product distribution. Furthermore, the styrene concentration at the electrode surface is controlled by different adsorption equilibria dependent on the electrode material (C, PbO 2, and Pt) [85]. β-Alkyl substituents decrease the yield of dimers and favor the formation of dimethoxylated monomers. On the other hand, the Y

–e–

Y –e–

+2HNu, –2H+

a

+ Y

+

Y

Y

+

b

–2H+

Nu Y

Y Nu Y

Y

+ Y

Y

c +

HNu,–H

–e– +

Y HNu,–H+

Nu Nu Y: phenyl, vinyl, alkoxy, amino, alkyl; Nu: RO, RCO2

Y Nu

Nu

SCHEME 18.21 Pathways for the anodic coupling of olefins (Y = aryl, alkoxy, N-acylamino, thioalkyl, vinyl): path a, radical cation–substrate coupling; path b, radical cation–radical cation coupling; and path c, reaction of the radical cation with nucleophiles from the electrolyte.

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Oxidative Coupling

Soft graphite, MeOH NaOCH3, NaClO4,–e–

2

67% Soft graphite, MeOH a) Nal, NaOCH3 b) NaClO4, NaOCH3

CH3O

OCH3 +

–e– OCH3

SCHEME 18.22

OCH3 a) 20% b) 45%

a) 38% b) 20%

Anodic coupling of styrene and indene. 4-MeO-C6H4

C6H5

C6H5 38–26% A

H 79% A E pa: 1.14 V

64% A 1.68 V

67% A 1.59 V

20–45% B l–, CH3O–; ClO4–, CH3O– 1.56 V

a) Glassy carbon vs. Ag/AgCl A: Coupling to dimer, B: methanolysis to dimethoxylated monomer

SCHEME 18.23 Dependence of the coupling yield on the oxidation potential and the β-substitution of the aryl olefin; the arrows indicate the coupling site; %, yield; Ep, oxidation potential in cyclic voltammetry.

coupling is favored by hydrogen in the β-position and a better stabilization of the positive charge in the radical cation, which is related to the oxidation potential of the olefin (Scheme 18.23). Reasons for the increased coupling would be a lower steric hindrance for the β-coupling and a better separation of the positive charge at the α-carbon and the spin at the β-carbon atom [12] (see also Section III.E.2). For the dimerization of 4,4′-dimethoxystilbene, it has been possible to demonstrate spectroelectrochemically [86] and at the rotating disc electrode [87] that the product is formed mainly by radical dimerization of the intermediate radical cations (Scheme 18.21, path b). Fast derivative cyclovoltammetry (CV), however, supports for the same olefin a complex electron transfer–chemical reaction–electron transfer (ECE) pathway (Scheme 18.21, path a) [88]. Depending on the kind of nucleophiles (acetate, water, or methanol) and the anode potential, one obtains a tetrahydronaphthalene derivative 44 by dimerization and electrophilic aromatic substitution (Scheme 18.24) [89], a monomer diacetate [89], or a 1,4-dimethoxy dimer [83]. When methanol is replaced by aqueous DCM or by an aqueous AN emulsions as solvent, styrene and α-methyl styrene yield 2,5-diphenyltetrahydrofuran (45) (Scheme 18.24) [90] and 2,5-diphenyl-2,5-dimethyl-tetrahydrofuran [91], respectively. OCH3

AcO

O

CH3O OCH3

OCH3

44

SCHEME 18.24 Products from anodic oxidation of styrene.

© 2016 by Taylor & Francis Group, LLC

45

720

Organic Electrochemistry

3. Dienes Butadiene is dimerized to dimethoxyoctadienes and forms trimers and dimethoxylated monomers as side products (Scheme 18.25) [92]. Dimerization of dienes is favored by soft graphite or carbon cloth anodes, by high olefin concentrations, and by terminally unsubstituted dienes (Table 18.1) [93,94]. Anodes with a smooth surface, such as platinum, gold, and glassy carbon, promote the formation of monomers; porous materials, such as soft graphite, favor the dimers. Qualitatively, the result can be explained by lower current densities due to the larger real surface of the electrode; this favors the coupling of the radical cation with the substrate (Scheme 18.21, path a). On the other hand, dimethoxylated monomers are favored against dimethoxylated dimers by alkyl substituents in the terminal position that hinder the C–C bond formation of the radical cation. 1,4-Dimethoxy adducts are the major isomers in monomers formed from dienes; a 1,2-dimethoxy product is formed in the monomer from the triene: β-ionone (Table 18.1). 4. Enols and Enol Ethers Oxidation potentials of 40 enols, enolates, and some selected α-carbonyl radicals are presented along with their characterization by various techniques (X-ray, EPR, ENDOR, general TRIPLE Monomer OCH3

OCH3

+

+ CH3O

CH3O

CH3O

C 0% Pt 24%

OCH3

OCH3 +

MeOH, NaClO4

CH3O

CH3O

OCH3

Dimer +

C, Pt, –e–

+

+

OCH3 C 49% Pt 9%

OCH3 + 2 isomers Trimer

+ CH3O

+

C 10% Pt 0%

+

SCHEME 18.25 Anodic coupling of butadiene: C, soft graphite anode; Pt, smooth platinum anode; in the boxes, assumed cationic intermediates.

TAbLE 18.1 Dienes: Oxidation Potentials and Products of Anodic Oxidation Diene 1,3-Butadiene Isoprene 1,3-Pentadiene 2,4-Hexadiene 1,3-Cyclohexadiene β-Ionone a b

Ep (V)a

Monomer (%)

Dimer (%)

Trimer (%)

Oligomer (%)

2.0 1.75 1.48 1.28 1.36 1.12

0 — 24 54 39 73b

49 26 29 6 6 —

10 8 — — — —

— >30 — — — —

At a glassy carbon electrode, versus Ag/AgCl. 9,10-Dimethoxy-9,10-dihydro-β-ionone.

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Oxidative Coupling

magnetic susceptibility measurements, UV–vis, fast-scan cyclic voltammetry, isotope effects). The compounds comprise stable enols linked to a multitude of substituents (alkyl, alkenyl, alkynyl, aryl, heteroaryl, propargyl alcohol) and stable enols of amides. The results allow to clarify the primary reaction pathway of enol radical cations as a rapid deprotonation and—if permitted by the oxidation potential and the strength of the oxidant—a follow-up oxidation of the α-carbonyl radical to the α-carbonyl cation. Moreover, the experimental oxidation potentials were linearly correlated with AM1-computed ionization potentials after correction for solvation. The correlation allows a reliable prediction of oxidation potentials of radicals including α-carbonyl radicals [95]. Enol ethers can be easily prepared from carbonyl compounds. Their anodic oxidation in methanol, NaI or methanol, and NaClO4 provides in 30–60% yield acetals or ketals from 1,4-dicarbonyl compounds (Scheme 18.26) [92,96–98]. The dimer yield is much less susceptible to an alkyl group at the coupling site than the aryl olefins and dienes. Apparently the good stabilization of the positive charge by the alkoxy group allows a good separation of positive charge and spin density in the radical cation that favors radical coupling. The β,β′-coupling of the monomers and the electrochemical reaction order (υenolether = 1.0, υmethanol = 0) support the radical cation as intermediate in the dimerization (Scheme 18.27). Silyl enol ethers (Scheme 18.21, Y = OSi(CH3)3) can be dimerized to 1,4-dicarbonyl compounds in good yields (Scheme 18.28) [99]. This way, unsymmetrical ketones can be coupled selectively in the α- or α′-position, since the corresponding silyl enol ethers can be prepared regioselectively. To suppress the methanolysis of the silyl enol ethers, AN and 5% MeOH are used as solvents and the electrolysis is conducted within 1 h by the use of a capillary gap cell, which allows a fast conversion. Vinyl ethyl ether can be linked to styrene in a cross-coupling reaction (Scheme 18.29) [100]. Further cross-couplings were obtained with α-ethoxystyrene and vinyl ethyl ether (17%), α-methylstyrene and vinyl ethyl ether (32%), and styrene and butadiene (22%). 5. Enamines Dimethylamino-substituted alkenes are easily oxidized at the anode. Their oxidation potentials range from −0.9 to 0.7 V (vs. SCE). The cation radicals can have long lifetimes, and ESR studies show that O ROH, H+ R1

OR

R1

–2e–, –2H+ +2MeOH

R1 OMe

RO

OR

MeO

R1

SCHEME 18.26 Preparation of enol ethers and their anodic dimerization.

51% (E1/2 = 1.72 V, n = 2.2)

OEt

OMe

OEt 59% (1.44, 1.7)

61% (1.30, 1.1) OEt

OEt EtO

50% (1.25, 1.0)

48% (1.28, 1.15)

50% (1.27, 1.1)

SCHEME 18.27 Anodic coupling of enol ethers in β-position to the alkoxy group: yields, oxidation potentials (vs. Ag/AgCl), and n-values. (Adapted from Koch, D. et al., Chem. Ber., 107, 3640, 1974.)

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Organic Electrochemistry OTMS

OTMS 1.33 V, n = 1.1, 66%

1.3 V, n = 1.0, 58%

OTMS OTMS C4H9 1.42 V, n = 0.7, 58%

1.35 V, n = 1.2, 78%

TMSO 0.95 V, n = 1.0, Hydrolysis only

SCHEME 18.28 Anodic coupling of silyl enol ethers: oxidation potentials (vs. Ag/AgCl), n-values, and dimer yields. C2H5O

OCH3 +

OC2H5

–e–, C-anode MeOH, NaClO4

SCHEME 18.29

OCH3

38%

Cross-coupling between vinyl ethyl ether and styrene.

they are strongly polarized with only a small spin density at the dimethylamino group. The oxidation step and the fates of the radical cations are well characterized by cyclic voltammetry and ESR spectroscopy. However, preparative scale electrolyses with the isolation of products are not reported [101]. On the other hand, when vinylidinebisdialkylamines were oxidized by silver ion in AN, diamidinium salts were obtained as dimerization products [102]. Enaminoketones or enaminoesters yield via dimerization of the intermediate radical cations and subsequent ring closure pyrrole derivatives (Scheme 18.30) [103]. In 2-alkyl enaminoesters and amides, the deprotonation to pyrroles is blocked. Here 2,3-dibenzoyl-2,3-dialkyl-succinic dimethyl esters are isolated in moderate yield and high diastereoselectivity for the meso-product (Scheme 18.31) [104]. In the intermediate radical cation, charge and spin are well stabilized allowing the coupling to vicinal quaternary carbon atoms. The yields of dimer are decreased due to a competing hydrogen abstraction of the sterically shielded radical cation; thereby, the ketoester or ketoamide is formed that is used to prepare the enaminoester or amide, respectively. R1

H

–e–, NaClO4 NHR2

H3C R1

MeOH, graphite anode

R1

R1

R1 –2H+,–R2NH2

+

+ NH

N

NHR2

R2

R2 R1

R2 Yield (%)

CO2Me Bzl CO2Me H CO2Me CH3

SCHEME 18.30

45 34 36

Anodic coupling of enaminoesters to pyrrole derivatives.

© 2016 by Taylor & Francis Group, LLC

723

Oxidative Coupling O Bnz

H

N

Pt, NaClO4

Ph

MeOH, –e– OMe

Ph

O

R

O

O N

H

O

Pt, NaClO4 MeOH, –e–

OMe

R'

R' = –N

15%, 27%

O

Ph Ph

O

R

R O O R = Me, i–Pr: 36%, 35%, meso

R

Bnz

O OMe + OMe

O R' R'

O R'

+

O O 32% meso

27%

SCHEME 18.31 Anodic coupling of 2-alkyl enaminoesters and amides to meso-2,3-dibenzoyl-2,3-dialkylsuccinic dimethyl esters and amides.

NC CN

H2N

CN

–e–, MeOH, NaClO4 N H 21.5%

Ep = 1.25 V vs. SCE n = 1.26

CN

NC CN

H2N

–e–, MeOH, NaClO

4

N H 23%

Ep = 1.20 V vs. SCE n = 1.26 NC C6H4X H2N X H 4–CH3O

–e–, MeOH, NaClO CN

1.27 1.27

4

XC6H4

Ep (V vs. SCE) n 1.4 1.53

CN

N H

X H 4-CH3O

C6H4X Yield (%) 27 29

SCHEME 18.32 Anodic coupling of enaminonitriles to pyrrole derivatives in methanol and NaClO4: oxidation potentials, n-values, and yields.

Likewise enaminonitriles can be cyclodimerized by oxidation in MeOH and NaClO4 at a graphite electrode. The yields are moderate as the product is consumed due to further oxidation (Scheme 18.32) [105]. When the product precipitates from the electrolyte, the yield is much better (Scheme 18.33). The five-membered cyclic enaminonitrile affords the heterocyclic product 46 in 62% yield, whose low solubility prevents it from further oxidation. The reaction possibly proceeds, via a radical cation, coupling of the two mesomeric forms leading to a dication, which undergoes methanolysis, intramolecular addition to the ketimine, and deprotonation to the product. With the six-membered cyclic enaminonitrile, no corresponding dimer is obtained [105]. The structure of the product 46 is confirmed by an x-ray structure [106]. Aryl enaminones (with aryl substituents: X = NO2, Br, H, Me, and MeO) were prepared from p-substituted acetophenones. Their preparative oxidation in a divided cell at a Pt anode in MeOH and LiClO4 yielded the dimer and 2,5-bis-aryl-furane (Scheme 18.34) [107].

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Organic Electrochemistry

NC–(CH2)4–CN

NH2

MeOH, NaClO4

CN

Na, THF

N

CN

–e–, graphite NH2

N H

46

–e–

OCH3 61%

CH3OH, –2H+

CN NH2

+ NH2

N

+

CN +

H2N C

HOCH3

N

+ NH2

SCHEME 18.33

Anodic coupling of a cyclic enaminonitrile to the heterocyclic compound 46. Ar

O

O –e–, 1.2 V vs. SCE Ar

N

Ar =

SCHEME 18.34

MeOH, LiClO4

N O

NO2

Ar

N+ + Ar

O

50%

20%

Ar

Anodic coupling of aryl enaminones.

Enamines were oxidized in alkaline medium in the presence of anions of 1,3-dicarbonyl compounds derived from methyl acetoacetate, acetylacetone, and dimethyl malonate, to afford coupling products. Presumably the enamines are oxidized at the platinum anode to their radical cations at a lower potential than that of the anions. The radical cations react then with the anionic substrates to form substitution products by addition–deprotonation (Scheme 18.35) [108]. Chiral secondary amines react with aldehydes to form transient enamines that were oxidized with ceric ammonium nitrate (CAN) to electrophilic radical cations. These react with silyl enol ethers to yield α-substituted aldehydes with high levels of asymmetric induction. A wide choice of both the aldehyde and the enolsilane component can be applied giving access to a diverse assortment of enantioenriched 1,4-dicarbonyl compounds [109]. O

O N

MeOH, NaOMe, –e– + RH

N

Pt, 0.4–0.9 V vs. SCE

R

+ Double bond isomer R

Yield

60% CH(COCH3)2 CH(COCH3)CO2CH3 67% CH(CO2CH3)2 61%

SCHEME 18.35

Anodic addition of organic anions to enamines.

© 2016 by Taylor & Francis Group, LLC

725

Oxidative Coupling O O

H

–e–,

N H Bu4NClO4 (TBAP), DCM

+ H

N

O

N

O 88% O

Ph

Ph H O

Ph

N H

+

H

N

OTMS

O

Ph

N

TPAB, DCM,–e–

O 57% yield 64% ee

SCHEME 18.36 sec-amines.

α-Oxyamination of aldehydes with TEMPO by anodic oxidation in the presence of

In an undivided cell, TEMPO was oxidized to TEMPO+ in the presence of aldehydes and substoichiometric amounts of secondary amines to yield an aldehyde with TEMPO as α-substituent in good yield. The asymmetric variant of this α-oxyamination was examined using chiral secondary amines. It was confirmed by cyclic voltammetry that the radical cation of the enamine, arising from the sec-amine and the aldehyde, is the intermediate (Scheme 18.36) [110]. The enantioselective organocatalyzed α-alkylation of aldehydes with xanthene was achieved by anodic oxidation in the presence of chiral cyclic amines. The best yield (80%) and highest enantioselectivity (70% ee) were obtained with the chiral catalyst 47. On the basis of CV and DFT calculations, the authors propose the coupling of a xanthene radical with an enamine radical cation; both are generated at the anode (Scheme 18.37) [111]. 6. Alkenes with Alkyl Substituents Alkenes without electron-donating substituents are oxidized at potentials >1.6 V versus SCE and undergo allylic substitution by the solvent and formation of ketones by rearrangement [112]. Coupling becomes possible, when an unsymmetrical intermediate radical cation with a tertiary cationic site is involved (Scheme 18.38) [113]. O H O

–e–, Pt anode

+ H

Ph

DCM, TBAP

Ph

O

O

CH3

O N Bn

SCHEME 18.37

t-Bu N H 47

Organocatalyzed anodic α-alkylation of aldehydes with xanthene.

© 2016 by Taylor & Francis Group, LLC

726

Organic Electrochemistry 0.5 m NaClO4 MeOH, 1.6 V –e–

OMe + 6%

OMe 4% OMe

OMe + OMe

MeO

8%

SCHEME 18.38

OMe 59%

Coupling in the anodic oxidation of 2-ethyl-1-butene.

Ar

Ar

–e–, H2O, –H+

HO Ar

HO

Ar Oxdn.

Ar

50

–e–, H2O, –H+

Ar

49

48

2

Ar

OH

HO

Ar

Oxdn.

O

Ar

51

2

Ar Ar

O Ar

O

52

SCHEME 18.39 Anodic coupling of 1,2-diarylethyne.

7. Alkynes The activated alkyne, diarylethyne (48), has been coupled at the anode [114]. When ethyne 48 is oxidized in AN and LiClO4 at a graphite anode, the dimer 51 and the benzil (52) are obtained as main products (Scheme 18.39). For the product formation, the following pathway is assumed: The  radical cation reacts with residual water to the enol radical 49 that dimerizes to 50. Further oxidation converts 50 to 1,2-diaroyl-stilbene (51) (Ar = Ph: 17%; Ar = 4-CH3OC6H4: 47%). Further oxidation and hydrolysis of 49 leads to benzil (52) (Ar = Ph: 20%; Ar = 4-CH3OC6H4: 12%). According to a literature search, since this publication, no further examples of this interesting reaction have been reported. 8. Alkanes The dehydrodimerization of short-chain hydrocarbons has been achieved in some examples. Coupling of methane at a Ag/Bi2O3 anode in a solid electrolyte yielded at 700°C and 2% conversion in 72% selectivity ethane and in 18% selectivity ethene [115]. Further dehydrodimerizations of methane are reported at 500–900°C for a Y2O3/ZrO2 solid electrolyte and different anode materials like Ag, Ni, Cu, Bi, Pt, Sm, and Mn [116,117].

E.

INTRAMOLECULAR COUPLING

1. Aromatic Compounds A large number of 1,n-bis(methoxyphenyl)alkanes have been coupled intramolecularly to afford biphenyls being integrated into rings of different size. The method allows preparing in few steps meta-cyclophanes, polycyclic hydroaromatic compounds, and intermediates for the synthesis of steroids, heterocyclic compounds, and alkaloids. The section is organized according to an increasing number n of methylene groups, which in some cases are replaced by a heteroatom, mainly nitrogen.

© 2016 by Taylor & Francis Group, LLC

Oxidative Coupling

727

The coupling sites and the yields depend on the integer n, the position of the aryl substituent, the choice of the electrode–electrolyte system, the current density, and the oxidation potential. The coupling products are easily further oxidized to their radical cations. In AN these are fairly unstable, which causes lower yields in this solvent. In DCM and TFA, the cyclized products are also further oxidized, but in this solvent, the radical cations are stable. After the electrolysis, they are reduced with zinc dust to the neutral products that can then be isolated in high yield [118]. Two examples are shown in Table 18.2, entry 1 and 2. 1,n-Bis(methoxyphenyl)alkanes with n = 1–4 and without a heteroatom in the alkyl chain cyclize in moderate to high yield (Table 18.2, entries 1–7, 11, 12). CV and products of unsymmetrically methoxylated diphenylalkanes permit conclusions on the mechanism of the dimerization. The results point to the dimerization of two radical cations [122]. On the other hand, kinetic studies indicate a cyclization by a radical cation–substrate coupling, followed by a rate-determining electron transfer [127]. Diarylalkanes with a phenol ether group at one end and a phenolic group at the other end afford high yields of spirodienones (Table 18.2, entries 3 and 4). The 2-methyl-tetramethoxybibenzyl derivative 53 cyclizes in an anodic 2,6′-coupling reaction followed by a dienone–phenol rearrangement to afford a 98% yield of the dihydrophenanthrene 54 (Scheme 18.40). Compound 54 contains essential structural elements of the B, C, and D ring of steroids [128]. With the corresponding dimethoxybibenzyl derivative, ring closure takes place without rearrangement to form a structure analogous to the A, B, and C ring of steroids [129]; however, the yield is moderate (22%) because of the lower activation of the arene system. A remote substituent can strongly influence the occurrence or absence of a further dienone–phenol rearrangement (Table 18.2, entries 6 and 7). In the case of the [2.2]-meta-cyclophane 55, two anisyl radical cations couple in 90% yield to give the tetrahydropyrene derivative 56 (Scheme 18.41) [130]. [2.2]-Meta-cyclophanes with other substituents than methoxy—for example, with CH3, Br, CN, or NO2—can also be cyclized [131]. The anodic cyclization of benzyltetrahydroisoquinolines 57 (Scheme 18.42, Table 18.3) with n = 2 and no heteroatom in the methylene chain to the morphinane skeleton is particularly interesting from the preparative point of view, since the cyclization cannot be achieved so simply with chemical oxidants. Compounds 58a,b can be obtained in good yield in AN and sodium bicarbonate using a divided cell [129,132–135]. Flavinantines 58c–f can even be prepared in an undivided cell and an acidic electrolyte [136]. The reaction has also been used to prepare optically active 9-(R)-Omethylflavinantine [137]. However, in order to obtain the desired morphine-type compound, one needs a C2′–C4a coupling in 57, but the −I effect and the steric shielding by the 3′-methoxy group favors the unwanted C6′–C4a coupling in 57 to the flavinantine type of compound. Attempts to enforce the coupling to the morphine type compound have so far been partially successful. Compound 59a (Scheme 18.43, Table 18.4) cyclizes at C2′ with debromination (due to R1 = Br, the numbering has changed in the benzyl ring compared to 57) [135]. In compounds 59b [129,135] and 59c,d [138], oxidative cleavage of the molecule is occurring at C1. In compounds 59e,f, the C2′ ring closure is not hindered by the iodo or bromo atom at C3′ [129,135]. In compound 59g, however, where the C2′ and C6′ coupling is identical, a C–C bond ortho to the methoxy group is obtained in 35–55% yield [133,135]. Compound 59h, which also has two identical ortho-coupling sites could be cyclized in 65% yield, and the cyclization product subsequently has been converted into 3-desoxythebain [139]. Benzyltetrahydroisoquinoline 60 with a differently substituted 3′- and 5′-position in the benzyl ring could be cyclized in 65% yield to compound 61, which was further converted to the thebaine derivative 62 (Scheme 18.44) [140]. A major mechanistic study, using mainly CV, and seeking common denominators in the oxidation of methoxybibenzyls including benzyltetrahydroisoquinolines, has been published [134].

© 2016 by Taylor & Francis Group, LLC

728

Organic Electrochemistry

TAbLE 18.2 Intramolecular Coupling of Aryl Compounds and Tin-Compounds Entry 1

Substrate

Condition

Product

DCM, TFA, TBABF4, Pt

OMe OMe

OMe

DCM, TFA, TBABF4, Pt

OMe

OMe

[119]

85

[120, 121]

75

[120, 121]

33

[122]

75

[123]

OMe

OMe OMe AN, MeO TMABF4, Pt 0.74V(SCE) MeO

MeO MeO MeO

OMe

MeO

OMe

O

OH Pt, AN, NaClO4

OH

O

O

O

O

O Pt, DCM, TFA (3:1) TBABF4

MeO MeO

MeO MeO

MeO

MeO

MeO

MeO 6

95 OMe

OMe

5

[119]

OMe

OMe

4

95

OMe

OMe

3

References

OMe

OMe

2

yield (%)

NRMe MeO MeO

AN, HBF4, Pt, R = CO2Et OMe 1.0 V(SCE)

NRMe MeO O

OMe OMe OMe (Continued)

© 2016 by Taylor & Francis Group, LLC

729

Oxidative Coupling

TAbLE 18.2 (Continued) Intramolecular Coupling of Aryl Compounds and Tin-Compounds Entry

Substrate

7

Condition AN, HBF4 Pt, R = COCF3

NRMe MeO

Product

yield (%)

OMe

References

75

[123]

45

[124]

43

[124]

62

[143]

36

[125]

MeO

OMe

MeO

OMe NRMe

MeO O 8

NaOH, NaClO4, C

O

O O

O NH

O 9

N

O NaOH, NaClO4, C

O MeO

O MeO NH

NH MeO

MeO 10

C, AN NaClO4

O n

O H

O

N H MeO

N

OMe

O 11

H2O, C Et4NClO4

MeO NCH3

HO

MeO NCH3

HO

O OMe MeO OH 12

MeO MeO

N

CH3

OMe OMe

AN, NaClO4 1.2 V(SCE)

[126]

MeO MeO

N +

CH3

OMe OMe (Continued)

© 2016 by Taylor & Francis Group, LLC

730

Organic Electrochemistry

TAbLE 18.2 (Continued) Intramolecular Coupling of Aryl Compounds and Tin-Compounds Entry

Substrate

13

Condition

Y R

Product

DCM, Bu4NBF4 molecular sieve, C

SnBu3

n

yield (%) 50–98

Y R

n

References [146]

F

R = C7H15,C6H12 Y = O, NCO2Me n = 1,2 14

DCM, Bu4NClO4

SnBu3 Y

[147]

Y R

R

Y = O, R = H Y = O, R = F Y = NCO2Me, R=H

OMe

OMe

OMe

OMe +

6'

MeO

MeO

–2e–

2

MeO

55 80 54

+ MeO

1 53

OMe

Rearrangement –2H+

O MeO OMe 54

SCHEME 18.40 Anodic 2,6′-coupling of the 2-methyl-tetramethoxybibenzyl derivative 53 to the dihydrophenanthrene 54.

OMe

MeO

–2e–, –2H+ 90%

OMe

MeO

55

SCHEME 18.41

56

Anodic coupling of the [2.2]-meta-cyclophane 55 to the tetrahydropyrene derivative 56. OR3 R 1O R 2O

R4O

4a NMe 2´ 6'

57

OR3 OR4

–e– AN,NaHCO3 or AN, HBF4

NMe R1O

58 O

SCHEME 18.42

Anodic cyclization of benzyltetrahydroisoquinolines 57 to flavinantines 58.

© 2016 by Taylor & Francis Group, LLC

731

Oxidative Coupling

TAbLE 18.3 Structure and yields of Flavinantines 58 a b c d e f

R1

R2

R3

CH3 CH3 CH3 CH3 CH3 CH3

CH3 CH3 CH3 CH3 CH3 CH3

R4

yield (%)

CH3 CH3 CH3 Bn CH3 H H CH3 –(CH2)– CH3 CH3

52–85 43–63 63 50 80 70

R4 R3

6' 2'

R2

R1

1 NR5

MeO OMe

59

SCHEME 18.43 compounds.

Attempts to enforce the coupling of benyltetrahydroisoquinolines 59 to morphine-type

TAbLE 18.4 Substituents in 59 a b c d e f g h

R1

R2

R3

R4

R5

Br Cl NO2 NHCOCH3 H H H H

H H H H I Br MeO BnO

MeO MeO MeO MeO MeO MeO MeO H

MeO MeO MeO MeO MeO MeO MeO BnO

Me Me Me Me Me Me Me COCF3

Benzyl- and 2-phenethyl-enaminoketones can be cyclized in good yields to quinolines and benzazepines (Table 18.2, entries 8 and 9). The cyclization has been also used to prepare a precursor of the alkaloid lycorane (Scheme 18.45) [141]. In indoles, the oxidation of the heterocycle must be suppressed by electron-withdrawing substituents, which is the case in 3-keto-substituted indoles. α-Phenacetyl-indole (Table 18.2, entry 10, n = 1) undergoes no cyclization, which has been attributed to a geometrically unfavorable transition state [142]; on the other hand, the β-phenylpropionyl indole (Table 18.2, entry 10, n = 2) cyclizes in 62% yield [143]. The radical cation first couples to a spiro-intermediate, which then undergoes oxidation, bond breaking, rearrangement, and tautomerization to the indolenine. Unsymmetrical 1,3-diaryl propanes with phenolic hydroxy groups react to spirodienones, presumably via a phenoxonium ion (Table 18.2, entry 3). Trimethoxy-substituted diphenylpropanes form different products depending on the reaction conditions. In MeCN and HBF4 a spiro dienone is obtained (Scheme 18.46) [122]; in DCM and TFA (3:1), a seven-membered ring

© 2016 by Taylor & Francis Group, LLC

732

Organic Electrochemistry O

O

O

5'

O 3'

1. 0.01 M HBF4, AN –20°C, 2.5 mAcm–2

BnO

BnO

–e–

NCH3

NCH3 CH3O

CH3O

O 61

OCH3 60 O O 2. a) NaBH4, MeOH, 99%; b) Pd/C, 1,4-cyclohexadiene, EtOH, 91%; c) (i-Pr)2CHN(CH3)2, DCM, 68%

O NCH3 CH3O

62

SCHEME 18.44 Cyclization of the benzyltetrahydroisoquinoline 60 to compound 61 that was further converted into thebaine 62.

O

N

–e–, 38% MeOH

O

O

N

O O

SCHEME 18.45

O

Cyclization of a benzylenaminoketone to a precursor of the alkaloid lycorane.

OMe

OMe MeO

MeO AN, HBF4 –e–, 90%

O

MeO

DCM, TFA –e– OMe

MeO

MeO

MeO 71% +

MeO

MeO

SCHEME 18.46 Trimethoxy-substituted 1,3-diphenylpropane forms different products depending on the reaction conditions.

© 2016 by Taylor & Francis Group, LLC

733

Oxidative Coupling MeO R2

OAc R1

MeO MeO OMe 64

OH 63

OMe

O

62%

O

N COCF3

OMe

80% BnO

65

NCOCF3 Br

MeO 66

OMe

OMe MeO

OMe

MeO

OMe (CH2)n (CH2)n

MeO

MeO OMe 67

SCHEME 18.47

OMe OMe 68

OMe

Substrates and products from various cyclizations of aryl compounds.

is formed. The latter result is explained by the low nucleophilicity of the medium and catalysis of the dienone–phenol-like rearrangement by TFA. A 2-phenethyltetrahydroisoquinoline with hydroxyl groups as substituents leads to an ortho, para-coupled spiro compound (Table 18.2, entry 11); with alkoxy groups, a tetracyclic compound by nitrogen-aryl coupling is obtained (Table 18.2, entry 12). The intramolecular coupling of 63 (Scheme 18.47) has been used in a synthesis of colchicines [144]. Unsaturated enolacetates 64 (n = 3 and no heteroatom) have been cyclized to cyclohexenyl ketones [145]. With the stannyl group as electroauxiliary, oxonium or N-acylium ions can be generated, which react in an electrophilic intramolecular addition or substitution to afford cyclic six-membered ethers or amines or their benzo analogues (Table 18.2, entries 13 and 14). In the intramolecular coupling between phenols and olefins, Swenton has explored the dependence of the yield on the substituents in the aryl ring and in the olefin (Scheme 18.48) [49,148,149]. In these cyclizations, the intermediate phenoxonium ion adds to the double bond and the resulting cation is subsequently trapped by methanol. 1,4-Di(methoxyphenyl)butane (n = 4) forms in DCM and TFA initially a seven-membered ring followed by rearrangement to the eight-membered ring (Table 18.2, entry 5). The amide 65 in Scheme 18.47 (n = 3 and with the heteroatom substituted by the trifluoroacetyl group) has been cyclized to a precursor of (+)-oxocrinine [150]. Amide 66 has been coupled intramolecularly to a precursor of galanthamine [151]. 1,5-Bis(dimethoxyphenyl)pentane (n = 5) forms a seven-membered ring compound 67 by an aryl–benzyl coupling [119]. With bis(dimethoxyphenyl)alkanes, (n > 6) coupling occurs intermolecularly and with partial subsequent cyclization to form up to 40-membered rings 68 [152].

© 2016 by Taylor & Francis Group, LLC

734

Organic Electrochemistry R1

R1

R2 –2e–, –H+

+

O

HO

R2

Pt, –e–, AN/MeOH/ AcOH, LiClO4 R1

R2

R1

R2

OMe MeOH, –H+ O

O

R1

R2 H H CH3 H Ph H CH3 H –(CH2)3–

SCHEME 18.48 and olefins.

Yield(%) 16 85 65 35 69

Dependence of the yield on the substituents in the intramolecular coupling between phenols

2. Olefins The intramolecular anodic coupling of olefins derived from enol ethers, vinyl sulfides, and ketene acetals has been reviewed recently [4a]. Thereby, radical cations are involved that react with carbon, oxygen, and nitrogen nucleophiles as trapping groups. Intramolecular coupling is used to construct five-membered rings that could be applied for the synthesis of tricyclopentanoid ring systems, which are core structures in some biological active natural products. Thereby, an enol ether radical cation can be trapped intramolecularly by an allylsilane. The results show that the reaction is compatible with an allylic alkoxy group that is frequently needed in the target molecule. In addition, the reactions using either a trisubstituted allylsilane or a cis-disubstituted allylsilane group as the terminating group were found to lead to the formation of five-membered rings with kinetic control of the relative stereochemistry. The stereoselectivity of these reactions originated from an A1,3-interaction (allylic strain: 1,3-interaction between substituents at C1 and C3 of an allyl group that favors one conformer [153]) with the allyloxy group that raised the energy of the transition state leading to the cis-product. When the size of R in the allyloxy group was increased from t-butyldimethylsilyl (TBDMS) to triisopropylsilyl (TIPS), the stereoselectivity was further improved (Scheme 18.49) [154]. TMS RO

C anode, –e– Pt cathode

RO

RO +

LiClO4, MeOH/ THF, 2,6-lutidine OMe R = TBDMS R = TIPS

OMe

OMe OMe 67% 81%

OMe 8% —

SCHEME 18.49 Cyclization between an enol ether and an allyl silane as trap; kinetic control of the stereochemistry by the allylic alkoxy group.

© 2016 by Taylor & Francis Group, LLC

735

Oxidative Coupling H

H

O

Me

Me H

RVC anode C cathode H R

H Me

O X

AN, 10% MeOH, LiClO4, 2,6-lutidine

Y

O OMe

+

OMe

O

O Ot-Bu

OMe

OMe

MeO — 73% (X = OMe, Y = Me) 38% (X = Y = OC(O)O) —

R = OMe R = Me R = OC(O)Ot-Bu R = OTf

— — 26% —

SCHEME 18.50 A polarized radical cation favors cyclization by C–C bond formation (RVC = reticulated vitreous carbon).

The ability of the intramolecular olefin coupling at the anode to form new C–C bonds depends on the polarization of the intermediate radical cation rather than on how electron-rich it is. Substrates were studied that allowed for a direct comparison of these two parameters, namely, NMR data and oxidation potential. The successful cyclizations gave highly functionalized bicyclic molecules containing four contiguous stereogenic atoms, one of which was tetrasubstituted. Additionally, it was shown for the first time that an enediol ether derivative with an alkyl group and a t-butyloxycarbonyl group is compatible with the cyclization reaction. The cyclization tolerates a second donor group on the olefin leading to the radical cation, but only if such a group increases or maintains the polarization of the radical cation. Hence, for a ketene acetal–type substrate, it is beneficial to add the second donor group, but for an enediol-type substrate, it is important to make sure that the second oxygen substituent on the initiating olefin is electronically neutral. That means a more polarized radical cation favors C–C bond formation, while a less polarized radical cation supports a C–heteroatom bond formation (Scheme 18.50) [155]. The importance of polarized bonds in radical cations, when discussing their chemistry, has been also pointed out for aromatic compounds earlier [12]. The intramolecular anodic coupling of a ketene dithioacetal radical cation and an amide trapping group leads to a furanone; the coupling benefits greatly from the addition of water to the reaction medium. After optimization a yield of 83% has been obtained (Scheme 18.51) [156]. Intramolecular anodic olefin addition reactions can be used to synthesize furanose and pyranose C-glycosides [157]. Similarly, carbon–nitrogen bonds can be created to prepare pyrrolidine and piperidine rings [158] and proline derivatives [159]. The key step in the asymmetric synthesis of the sesquiterpene lactone (−)-alliacol A is the intramolecular anodic coupling of a silyl enol ether and a furan (Scheme 18.52) [160]. The electrolysis could be accomplished with a simple, cheap, and readily available equipment, namely, a threenecked round-bottom flask, graphite electrodes, and the use of a 6 V lantern battery [161]. The anodic cyclization to a tetrahydrofuran by trapping the radical cation of a ketene dithioacetal by an alcohol was used as a step in the asymmetric synthesis of (+)-nemorensic acid. Thereby, steric factors of the large ketene dithioacetal group led to high levels of stereoselectivity [162]. An anodic cyclization reaction between an enol ether radical cation and an oxygen nucleophile has been used to prepare a tetrahydropyran building block for the C10–C16 fragment of the biologically S

Et N Et

O

S

RVC anode, –e–, C cathode 10% H2O, MeOH O Et4NOTos, 2,6-lutidine

S

O

S OMe 83%

SCHEME 18.51

C–O bond formation by trapping a ketene dithioacetal radical cation with an amide group.

© 2016 by Taylor & Francis Group, LLC

736

Organic Electrochemistry TBSO

CH3

CH3

O a) RVC anode, –e– LiClO4, 2,6-lutidine,

TBSO

O

20% MeOH/DCM b) TsOH

HO

O 88%

SCHEME 18.52 Anodic intramolecular coupling of a silyl enol ether as key step in the asymmetric synthesis of (–)-alliacol A.

S

S

HO MeO

RVC anode, –e– Et4NOTs 30% MeOH/THF 2,6-lutidine, 45°C MeO OBn

S

S

MeO

O

HO

MeO

OBn

MeO

S

S

OBn 10%

70%

+S

OMe

+

S

+ HO

HO MeO

SCHEME 18.53

OBn

MeO

OBn

Product formation via an intramolecular electron transfer.

active, natural product bryostatin. The oxidative cyclization was successful despite the presence of a thioacetal group that has a lower oxidation potential than the enol ether. Experimental evidence suggested that the reaction proceeded through an initial oxidation of the thioacetal followed by an intramolecular electron transfer to form the enol ether radical cation that was subsequently trapped by the oxygen nucleophile (Scheme 18.53). The formation of the desired cyclic product could be explained using the Curtin–Hammett principle [163]. The scope of Curtin–Hammett-controlled anodic cyclizations has been further extended by examining the compatibility of coupling reactions conducted in the presence of a dithioketal protecting group with both the formation of medium-size rings and the generation of carbon–carbon bonds. Cyclizations utilizing an alcohol as trapping group are fast and show no signs of competitive trapping of the dithioketal radical cation even when forming a seven-membered ring. Carbon–carbon bond–forming reactions utilizing reactive furan trapping groups are also fast, although not fast enough to allow for seven-membered ring formation. Slower carbon–carbon bond–forming reactions benefit strongly from the use of less polar electrolytes that selectively lower the concentration of methanol in the region of the reaction at the anode surface [164]. Intramolecular coupling reactions between enol ether radical cations and oxygen nucleophiles are also governed by stereoelectronics. By taking advantage of this observation, a THF building block for use in constructing (+)-linalool oxide and rotundisine has been synthesized in four steps from a commercially available starting material [165]. Intramolecular anodic coupling of two enol ethers can be used to build bicyclo[3.2.1]octane ring skeletons. Wider applications are limited due the formation of dimethoxy acetals at both the terminating and initiating ends of the cyclization reactions. However, these ends can be differentiated by using a ketene dithioacetal group and an enol ether as initiating and terminating group (Scheme 18.54) [166].

© 2016 by Taylor & Francis Group, LLC

737

Oxidative Coupling

SCHEME 18.54

S

RVC anode, –e– Pt cathode

S

30% MeOH/THF 2,6-lutidine

OMe

S S MeO OMe MeO 75%

Intramolecular coupling between an enol ether and a ketene dithioacetal.

R1

S S

NH2 R3

n

R2

RVC anode, –e– Pt cathode LiOMe MeOH, Et4NOTs

R3

H S N

S OMe R1

n

R2

R1 = Me, CH = CHMe; R2 = H, Me; R3 = H, Me Yield: 72–92%, high diastereoselectivity

SCHEME 18.55 Intramolecular trapping of a dithioketene acetal by an unprotected amine to generate amino acid derivatives.

Anodic oxidation of enol ethers, for example, HOCH2(CH2)nCHRCH=CHOMe (n = 1, 2), was used to accomplish the formation of both tetrahydrofurans and tetrahydropyrans. The method compliments existing chemical routes to related ring systems [167]. Inexpensive, readily available photovoltaic cells can been used to conduct electrochemical oxidation reactions. It has been applied to intramolecular coupling reactions between two enol ethers and between enol ethers and oxygen or nitrogen nucleophiles. In this way, the energy efficiency of sunlight-driven reactions can be combined with the versatility of electrochemistry to create new, sustainable methods for conducting oxidations [168]. A radical cation generated from a dithioketene acetal can be trapped by an unprotected amine. The reaction generates amino acid derivatives with a tetrasubstituted α-carbon atom. In this reaction the cyclization lowers the oxidation potential of the substrate. In the examples reported, the cyclizations are fast enough that the substrate potential is decreased to such an extent that it is significantly lower than that of the product, which this way can be prevented from further oxidation. The thioketal can be hydrolyzed with N-chlorosuccinimide (NCS) in acetone and water to unmask the amino acid derivative (Scheme 18.55) [169]. The anodic oxidation of bis-enol ethers can effectively lead to the formation of five-, six-, and seven-membered ring products in good yields. Besides leading to potentially useful 1,4-dicarbonyl equivalents, the cyclizations are compatible with the formation of quaternary carbon atoms [170]. Enol ethers could be coupled with aryl rings to form a fused six-membered ring in 57% yield. The cyclization profited from a 3-alkoxy group in the aryl ring. Furan rings were found to be excellent coupling partners for the reactions and afforded products having fused, bicyclic furan ring skeletons. Cyclizations involving furans were shown to be compatible with the formation of both six- and seven-membered rings, the generation of quaternary carbon atoms, and the use of a variety of electron-rich olefins as the other coupling partner. The yields ranged between 10% and 71%. It appeared that the furan can serve as either the initiating group or the terminating group for the cyclizations. Also N-acylpyrroles could be coupled to thioenol ethers in 66% yield [171]. Two enol ethers or an enol ether and an allylsilane could be coupled to form fused bicyclic compounds consisting of five- and six-membered rings. The resulting 1,4-dicarbonyl compounds could be subsequently used for aldol reactions to form angularly fused tricyclic ring systems [172]. Enol ethers could be coupled intramolecularly to simple alkyl olefins, styrenes,

© 2016 by Taylor & Francis Group, LLC

738

Organic Electrochemistry

O

TIPSO N

O

O

O

RVC anode and cathode 2,6-lutidine, MeOH/THF ET4NOTos, –e–

N

O

n

n

OMe OMe

OMe

n = 1 87% n = 2 65% O

TIPSO N n

CH2TMS

SCHEME 18.56 allylsilane.

O

RVC anode and cathode 2,6-lutidine, MeOH/THF ET4NOTos, –e–

O

O N

O

n

n = 1 70% n = 2 71%

Intramolecular coupling reactions of N,O-ketene acetals with enol ethers and with an

and allylsilanes in isolated yields ranging from 57 to 84%. The reactions were found to be effective for generating both five- and six-membered rings [173]. Intramolecular coupling reactions of N,O-ketene acetals with both enol ether and allylsilane terminating groups were examined. The reactions of the N,O-ketene acetals with allylsilane groups were found to be much more efficient than corresponding reactions with dithioketene acetal groups and allylsilanes. The reactions were also more efficient than the intramolecular coupling reactions between enol ethers and allylsilanes (Scheme 18.56) [174]. Both results support the observation that more polar radical cations aid the formation of carbon–carbon bonds. A competition experiment was designed where an N-nucleophile and an O-nucleophile compete in the coupling reaction. The studies show that when a dithioketene acetal olefin is used in the reaction, the use of LiOMe as a base leads to a mechanism, where probably the radical cation is trapped by a sulfonamide anion. A similar result was obtained when an enol ether was used. However, the data obtained with a vinyl sulfide were less conclusive [175]. Two intramolecular anodic olefin coupling reactions have been used to synthesize the arteannuin ring skeleton. In one case an N,O-ketene acetal was the initiating group and a furan terminated the reaction. In the second case a (Z)-enolether initiated the coupling and installed a tetrasubstituted carbon atom (Scheme 18.57) [176]. A two-step electrochemical annulation has been developed for the preparation of fused furans. The process involves an initial conjugate addition of a furylethyl cuprate and trapping of the enolate as the corresponding silyl enol ether. The second step of the annulation involves the anodic coupling of the furan and the silyl enol ether to form a six-membered ring [177]. Cyclic voltammetry indicates that in the anodic oxidation, the enol ether is converted to a radical cation that is trapped by an electron-rich heterocyclic compound (furan, thiophene) with the formation of six- and seven-membered rings. The dramatic influence of the methyl group is attributed to a kind of gem-dialkyl effect. Mechanistic considerations are well supported by cyclic voltammetry, suitably chosen probes, and the use of competition experiments (Scheme 18.58) [178,179]. The effect of the methyl group was extended to other groups as phenyl, i-propyl, ethyl, and vinyl in this position that led to consistent good yields of around 60% of the seven-membered rings. Voltammetric studies indicated that this effect lowers the oxidation potential by approximately 110 mV.

© 2016 by Taylor & Francis Group, LLC

739

Oxidative Coupling

OTIPS CH3

O O

N O

Ph

O

O 1) RVC anode, –e– carbon cathode LiClO4, MeOH, DCM O 2,6-lutidine

N

2) H3O+ work up

Ph

R

RVC anode carbon cathode LiClO4, MeOH, DCM 2,6-lutidine

MeO O

O

R=H R = CH3

Me

H

CH3

H

70% 65%

R

Me

O O

R

R MeO R=H 70% R = CH3 80%

SCHEME 18.57

Two cyclization routes to the arteannuin skeleton. R

TMSO

R

RVC anode,–e–, LiClO4, AN, iPrOH, 2,6-lutidine O O

O

R=H R = CH3

SCHEME 18.58

0% 70%

Annulated heterocycles through a radical cation cyclization.

This result provides direct evidence for the importance of the reactive rotamer effect on an electron transfer reaction [180]. The shift of the oxidation potential could be due to an assistance of the electron transfer by the π-electrons of the furan ring. This would favor the coupling reaction against other competing reactions of the enol ether radical cation. A concise, stereoselective, and convergent total synthesis of the unnatural enantiomer of the neodolastane diterpenoid heptemerone B was completed. Thereby, the central seven-membered ring was closed by an electrochemical oxidation (Scheme 18.59) [181]. An alternative annulation of enol ethers located in five- and six-membered carbocyclic rings by cyclization with furans and thiophenes has been developed. The connection of the furan or thiophene ring at C2 to the enol ether led to spiro compounds (Scheme 18.60) [182]. Me

Me Me

TBDPSO O

OBn OTBS

Me

RVC anode, –e– 2,6-lutidine LiClO4, 20% MeOH, DCM

H TBDPSO MeO

O

OBn

H O

81%

SCHEME 18.59 Synthesis of the central 7-membered ring in the 17-step total synthesis of the diterpenoid heptemerone B using anodic coupling as one of the key steps.

© 2016 by Taylor & Francis Group, LLC

740

Organic Electrochemistry R

R

R C anode, –e– AN/iPrOH LiClO4 O 2,6-lutidine

n TMSO

n

n

+ O

O

O

O OiPr

iPrO 22–23%

59–69%

R = H, Me, n = 1, 2

SCHEME 18.60 Anodic annulation of five- and six-membered rings to furan leading to spiro compounds. OMe

OMe

MeO

C anode, –e– 0.5 mA cm–2

MeO

CH3

AN, LiClO4

CH3 67%

TMSO

SCHEME 18.61 skeleton.

O

Anodic coupling of a silyl enol ether to an electron-rich aryl ring as a way to the hamigeran

R

R OH N Ph

N OH

N Ph

N N

N

N Ph

O

MeOH, –e– NaBr O

N

Ph 84–97% R: H, 3-Br, 2-Cl, 4-Me, 2-MeO, 3-Cl, 4-Et

SCHEME 18.62 Mediated intramolecular coupling of heterocyclic compounds.

The tricyclic core of the hamigerans, which are unusual halogenated marine natural products, has been prepared through the use of a two-step electrochemical benzannulation reaction. The annulation proceeds through an initial conjugate addition of a phenethyl cuprate to 3-methylcyclopentenone with in situ silylation of the resulting enolate. Anodic oxidation couples the pendant arene and the silyl enol ether to produce a key intermediate for the hamigeran synthesis. Careful optimization revealed that the use of alcohol-free AN as solvent was critical to obtain high yields of the annulated product (Scheme 18.61) [183]. The intramolecular coupling of electron-rich heterocyclic compounds has been achieved in methanol with NaBr as mediator in high yields (Scheme 18.62) [184].

F. ANODICALLY INDUCED CYCLOADDITIONS 1. [2 + 2] Cycloadditions Vinylcarbazoles and 1,1-bis(4-dimethylaminophenyl)-ethylene can be converted with Fe(NO3)3 in methanol in 60–90% yield to the cyclodimers. When the vinyl group is attached to the N-atom, head-to-head cyclobutanes are obtained, and in case the vinyl group is attached to the aryl ring and in 1,1-diarylethylenes, head-to-tail cyclobutanes are formed. As for the mechanism, it is assumed

© 2016 by Taylor & Francis Group, LLC

741

Oxidative Coupling H H

CH3O –e– CH3O

SCHEME 18.63

CH3O

OCH3

CH3O

OCH3

H H 30%

Cycloaddition of dimethoxyindene to a cyclobutane. CH3

–e–, AN, Et4NClO4

Ph

+ CH3

O

O CH3 Ph

H3C Ph

CH3 O 15%

+

H Ph

O O

CH3

O

50%

SCHEME 18.64 [2 + 2] and [4 + 2] cycloadducts from dimethylindenone and 1-benzoyl-1-phenyl-ethene.

that the olefin is oxidized to a radical cation that couples with the olefin to a 1,4-radical cation that undergoes cyclization. The cyclic radical cation is reduced to the cyclobutane by an electron transfer from the olefin, which continues the chain reaction [185,186]. At the anode, N-vinylcarbazole has been cyclodimerized in AN even earlier, however, in only 8% yield [187]. In acetic acid KOAc as supporting electrolyte, a concentrated solution of dimethoxyindene is dimerized with 1.4 F in 30% yield in a formal [2 + 2] cycloaddition to a cyclodimer; polymers, however, are the major products (Scheme 18.63) [188]. Dimethylindenone and 1-benzoyl-1-phenyl-ethene combine, when electrolyzed in AN and Et4NClO4 to form 15% of a [2 + 2] cycloadduct and 50% of a [4 + 2] cycloadduct (Scheme 18.64) [189]. Cis-cyclooctene shows no voltammetric response in the electrolyte DCM and Bu4NB(C6H5)4 prior to the anodic background. However, in the presence of catalytic amounts of ReCp(CO)3 and its oxidation at 1.16 V versus FeCp2 to the strong oxidant [ReCp(CO)3]+, cis-cyclooctene is oxidized to form the cyclodimer in 70–80% yield. Presumably, [ReCp(CO)3]+ converts the cycloolefin to the radical cation that undergoes cyclodimerization with cis-cyclooctene as donor (Scheme 18.65) [190]. Cross-metathesis between an enol ether and an aryl alkene can be selectively initiated at the anode. Thereby, the 4-methoxy group at the aryl ring determines the pathway. The following mechanistic explanation is given: The enol ether is oxidized to a radical cation that undergoes a cycloaddition with the 4-phenyl-1-butene to form a cyclobutane radical cation. This radical cation forms in a retro-cycloaddition the cross-metathesis product and an enol ether radical cation, which carries on the chain (Scheme 18.66a). With the anisole group instead of the phenyl group, the electron is faster transferred from the electron-richer aryl ring to the cyclobutane radical cation, thus preventing the retro-cycloadditions and leading selectively to the cross-coupling product (Scheme 18.66b) [191]. –e–, ReCp(CO)3 DCM, Bu4NB(C6H5)4

70–80%

SCHEME 18.65

Anodic cyclodimerization of cis-cyclooctene mediated by ReCp(CO)3.

© 2016 by Taylor & Francis Group, LLC

742

Organic Electrochemistry

+ C8H17

–e–, 1.1 F

OCH3

LiClO4, MeNO2

(a) 20 mol equiv

78%

CH3O

CH3O + C8H17 (b)

C8H17

20 mol equiv

OCH3

OMe

–, 0.5 F

–e

LiClO4, MeNO2

+

C8H17

C8H17 3%

SCHEME 18.66 Selective formation of either a (a) cross-metathesis product or (b) a cross-coupling product controlled by the ease of the intramolecular electron transfer from the aryl group.

Further support on the reaction mechanism is given through deuterium-labeling studies [192]. The assumed intramolecular electron transfer from the alkoxyphenyl group was further supported by cyclic voltammetric studies and quantum chemical calculations [193]. Electron-transfer-induced intermolecular [2 + 2] cycloadditions between 3,4-dihydro-2H-pyran and several unactivated olefins have been demonstrated to work on the basis of the aromatic “redox tag” strategy. The anodically generated radical cation of the cyclic aliphatic enol ether was effectively trapped by unactivated olefins possessing an aromatic redox tag to stabilize the corresponding [2 + 2] cycloadducts. The aromatic redox tag is oxidized during the formation of the cyclobutane ring, affording a relatively long-lived aromatic radical cation, which is then reduced to complete the overall reaction. Aromatic tags with an oxidation potential between 1.46 and 1.76 V (vs. Ag/AgCl) most effectively increase the cyclobutane yield to 84–94%. Furthermore, the aromatic redox tag was also found to facilitate the electron-transfer-induced cycloreversion of the cyclobutane ring to the corresponding radical cations [194]. Anodic processes in the lithium perchlorate and nitromethane (LPC, NM) electrolyte system provide a wide variety of carbon–carbon bond–forming reactions. The electrolyte solution acts as a Lewis acid catalyst that enables intermolecular reactions between unactivated alkenes and anodically generated intermediates. The properties of the electrolyte solution and applications to electrochemical reactions are reviewed [195]. Alkyl enol ethers with a methoxyphenyl group attached to the alkyl group lead to high yields of cyclobutanes with methylene cyclohexane and other unactivated alkenes [196]. Further examples of electrocatalytic formal [2 + 2] cycloadditions between anodically activated aliphatic enol ethers and unactivated olefins possessing an alkoxyphenyl group were accomplished in an LPC and NM electrolyte solution [197]. Enyloxy benzene gave consistently high yields of cyclobutanes with enol ethers, an activated alkene and even an alkyne (Scheme 18.67) [198]. Formal hetero Diels–Alder reactions and [3 + 2] cycloadditions could be achieved by electrolysis of benzylic thioacetals in NM and LiClO4 in the presence of alkenes or allylsilanes. The reaction proceeded possibly by selective oxidation of the thiogroup to a radical cation and subsequent loss of a thioradical. Electrophilic addition of the benzyl cation to the alkene and subsequent aromatic substitution led to the product. With an o-hydroxy group in the aryl ring, an intermediate quinodimethane can be formed and react as a dienophile (Scheme 18.68) [199]. 2. [4 + 2] Cycloadditions At the anode [4 + 2] cycloadditions between two electron-rich dienes can be induced. Most probably, the diene is oxidized to the radical cation. This reacts with the dienophile to the cycloadduct radical cation, which is reduced to the product by the diene that thereby is transformed into the radical cation that continues the chain. 1,3-Cyclohexadiene forms in 24% yield the cycloadduct with a high endo/exo selectivity of 28:1. The yield, however, is much lower than in the photosensitized cycloaddition (95%, endo/exo = 8:1) (Scheme 18.69) [200].

© 2016 by Taylor & Francis Group, LLC

743

Oxidative Coupling O O

–e–, LiClO4, MeNO2 +

0.1 F 90%

O O

–e–, LiClO4, MeNO2 +

0.1 F O

O

96% O O

–e–, LiClO4, MeNO2 +

1.0 F 35%

SCHEME 18.67 The prop-1-enyloxy benzene group improves the yield in cycloadducts. OMe

OMe

OMe

–e–

–e– TMS

PhS

TMS LiClO4, MeNO2 –PhS

LiClO4, MeNO2 –PhS SPh PhS 100%

70%

–e–

OH PhS

SPh

LiClO4, MeNO2 –PhS

OCH3

O PhS 86% OCH3

–e–

OCH3

CH3O

OCH3

CH3O

LiClO4, MeNO2 –PhS PhS

SPh

PhS 79%

SCHEME 18.68 Hetero Diels–Alder reactions and [3 + 2] cycloadditions by anodic oxidation of benzyl thioacetals in the presence of alkenes.

2

AN, DCM, LiClO4 2,6-lutidine, –e–

SCHEME 18.69 Cycloaddition between two electron-rich dienes induced by single electron transfer (SET) at the anode.

© 2016 by Taylor & Francis Group, LLC

744

Organic Electrochemistry CO2CH3 HN CN

CN AN, LiClO4 –e–

N H

CO2CH3

N CO2CH3

HN

91%

HN –2H+, –e–

CN

CN + N H

N H + CO2CH3

HN

HN

SCHEME 18.70

[4 + 2] cycloaddition between an indole and an enaminoester.

[4 + 2] Cycloadditions between 2-vinylindoles acting as heterodienes and α-acceptor substituted cyclic and acyclic enamines can be induced by formation of 2-vinylindole radical cations either via anodic oxidation or photoelectron transfer (PET) using a catalytic amount of triarylpyrylium tetrafluoroborate as sensitizer. In this way, pyrido[1,2-a]indoles or indolo[1,2-a] hexahydro-1,8-naphthyridines are formed in one step with complete regiochemical and stereochemical control. The products formed are interesting as they incorporate the skeleton of indole alkaloids. Product formation occurs presumably by electrophilic addition of the radical cation to the enaminoester. This is followed by a ring closure through electrophilic attack of the carbocation at nitrogen. The reaction is finished by oxidation of the radical to a cation and deprotonation (Scheme 18.70) [201]. In other cases, formal cycloadditions are initiated by oxidation of phenols to quinones, which can react in Michael addition with enols (see Section III.B.2, Scheme 18.12). Diels–Alder reactions of in situ generated quinones with dienes are accomplished in excellent yield in an aqueous sodium dodecyl sulfate (SDS) solution by selective anodic oxidation on a glassy carbon electrode being modified by a cation-exchange resin. Under these conditions the Diels– Alder reaction is markedly accelerated. This procedure provided Diels–Alder adducts in high yield (84–100%) and current efficiency (88–96%) (Scheme 18.71) [202]. OH

OH HO

+ CO2Et

–e–, SDS PTFE-coated anode

OH

CO2Et 100% O

CO2C8H17 + OH

O

CO2C8H17

–e–, SDS PTFE-coated anode

O

H 95%

SCHEME 18.71 Diels–Alder reaction of anodically generated o- and p-quinones in micellar solution and a polytetrafluoroethylene (PTFE)-coated anode.

© 2016 by Taylor & Francis Group, LLC

745

Oxidative Coupling R1

R1

OMe

R2

OMe

R2

OMe

OMe +

–2e–, –H+ R3

R3 R

R

OH [5 + 2]

[3 + 2]

O

[2 + 2]

O

O

O

CH3O

MeO

R1

R

R3

O

R

R2

O

R2

R1

R

R1 R2

CH3O

R3

R3 O

SCHEME 18.72 Intramolecular coupling of a phenol derivative to [5 + 2], [3 + 2], and [2 + 2] cycloadducts.

The intramolecular coupling of phenol derivatives to a double bond can lead to three different product types [25]. The phenol is oxidized to a phenoxonium cation that can undergo either a [5 + 2] cycloaddition, a [3 + 2] cycloaddition, or a [2 + 2] cycloaddition (Scheme 18.72); see also Section III.B.2. The selectivity for the three pathways depends strongly on the configuration of the double bond and the groups R. The electrolyses are performed frequently in Ac2O and Bu4NBF4. Selective formation of the [5 + 2] adduct (59%) has been found with R = R2 = R3 = Me and R1 = H [203,204]. The [3 + 2] adduct (69%) has been obtained with the E-isomer and R = R1 = Me, R2 = H, and R3 = Ar [205]. With the same substituents but the Z-configuration of the double bond, the [2 + 2] cycloadduct (66%) was formed [206]. The [2 + 2] cycloadduct was also formed with R = Me, R1 = R2 = H, and R3 = Ar [207,208]. In a multistep one-pot electrochemical synthesis, a variety of 2-alkylamino-1,4-benzoxazine derivatives 73 are prepared, which are potential neuroprotective compounds (Scheme 18.73) [209]. Thereby, the monoanion of 1-benzoyl-3-amino-2,4-dihydroxybenzene (70) is oxidized at a Hg-anode at 0.05 V in MeOH and TEAP to the o-iminoquinone 69. The o-iminoquinone 69 is in equilibrium with 70 and works as a redox mediator for the indirect electrochemical oxidation of amines to the corresponding imines 71. The imines equilibrate with the enamines 72, which can OH

O

HN

R1 Ph

+ R3–NH2 NH2

R2

O

69 –e–

O

OH

H2N

R1

R1 Ph

R2

R2

NR3

–

O R2

OH

R1

HN

72

O

O

R1 Ph

+ R3NH

72

71

70

OH

H N

O Ph

R2 R3NH

69

NHR3

O

73

R1, R2 = Me, Me; Ph, Ph; cyclohexyl R3: (MeO)2CHCH2NH, cyclohexyl

SCHEME 18.73 Preparation of 1,4-benzoxazine derivatives 73 by an inverse Diels–Alder reaction between an in situ electrogenerated azaquinone 69 and an enamine 72.

© 2016 by Taylor & Francis Group, LLC

746

Organic Electrochemistry OH

O

HO

Ph +

–O

N O

Br OH

O

O

–e–, MeOH, Hg anode 0.05 V vs. SCE

Ph

O

N

Br

O 76%

SCHEME 18.74

Hetero Diels–Alder reaction of electrogenerated o-quinones with enamines.

undergo an inverse Diels–Alder reaction with the o-iminoquinone 69. When only one amine with R1 and R2 is used, the yields are up to 76%. With two amines—as shown—three groups R1, R2, and R3 are introduced with yields of 68–70%. To widen the molecular diversity, the o-iminoquinone 69 has been generated in the presence of different enamines [209]. The scope of the cascade reaction is extended by using substituted aryl rings in the benzoyl group and by replacing the benzoyl group by the cyclohexylcarbonyl, isobutylcarbonyl, and nitro group. Consistent high yields (71–80%) of 2-alkylaminobenzoxazines are obtained in the cascade reaction [210]. The anodic oxidation of pyrogallol derivatives produces chemically unstable o-quinones as heterodienes, which are trapped in situ by enamine dienophiles through inverse Diels–Alder reactions. The possibility of introducing variations in both cycloaddition partners gives rise to highly substituted 1,4-benzodioxin cycloadducts with up to five elements of diversity. The reactions proceed under mild conditions with good yields ranging between 25% and 77% as shown in 17 examples. The methodology should be amenable to the assembly of libraries of biologically relevant heterocycles (Scheme 18.74) [211]. The electrochemistry of selected benzoxazine derivatives has been studied, and furthermore, the compounds have been tested as neuroprotective agents for the treatment of cerebral palsy [212]. Anodic oxidation of appropriately substituted 2-methoxyphenols or (2-methoxyphenoxy)-2methylpropionic acids in the presence of methanol furnishes stable o-quinone monoketals in a CCE. The propionic acid derivatives are initially obtained as O-spirolactonic quinone bisketals that are then selectively hydrolyzed into the desired monoketals. In the absence of blocking substituents, o-quinone monoketals spontaneously undergo Diels–Alder dimerizations into tricyclododecadienedienones with high site selectivity, regioselectivity, and stereoselectivity (Scheme 18.75) [213]. O O

O

–e–, MeOH OCH3 LiCIO4, CCE

HO

O

O

O H+,H2O

O

OOMe OMe

O

[4 + 2]-CA

O

O

O

O

O

O

O 100%

O

SCHEME 18.75 Anodic oxidation of (2-methoxyphenoxy)-2-methylpropionic acid to a stable bisketal and the hydrolysis to a monoketal with a subsequent Diels–Alder dimerization.

© 2016 by Taylor & Francis Group, LLC

747

Oxidative Coupling Ph

Ph –e–, 0.5 eq. HCIO4

NC

N

O 0.5 F, 1.8 V vs. SCE

NC

n

OEt

+N

O n

Ph

O OEt

NC

N n

SCHEME 18.76

n = 1: 52% n = 2: 72% n = 3: 72%

Electrocatalytic cycloaddition of vinyl ethyl ether to an N-cyanomethyl-oxazolidine.

The anodic oxidation of N-cyanomethyloxazolidine initiates an electrocatalytic cycloaddition of vinyl ethyl ether at the N,O-acetal function, affording a two-carbon ring enlargement. As mechanism of the formation of a radical cation that undergoes ring opening to a distonic radical cation is postulated. Subsequent addition to vinyl ethyl ether, ring closure and 1e-reduction of the radical cation by the starting material yields both the product and a radical cation for continuation of the chain reaction. This assumption is supported by cyclic voltammetry and the understoichiometric consumption of electricity (Scheme 18.76) [214].

IV. COUPLINg VIA RADICALS A.

GENERAL COMMENTS

Radicals can be generated at the anode in two ways. In the first way, a C–H acidic compound is deprotonated to an anion that is oxidized. In preparative scale electrolyses, one uses mostly compounds with pKa values 250 mA cm−2) with current control in an undivided cell. High current densities favor coupling and increase the anode potential. For electrostatic reasons at a higher anode potential, the solvent methanol is displaced from the electrode surface and replaced by the negatively charged carboxylate. This way, the competing oxidation of methanol is suppressed. A supporting electrolyte is not needed as the carboxylate and the countercation act as supporting electrolyte. Foreign anions should be excluded because they can disturb the necessary formation of a carboxylate layer at the anode. Alkali and alkyl ammonium ions have no negative effect. In an undivided cell, the carboxylate concentration is held constant as, with each carboxylate ion that is converted at the anode, the equivalent of methoxide is formed by discharge of protons at the cathode. Thereby, the carboxylate is continuously regenerated from the excess acid. The endpoint of the electrolysis is reached, when the pH of the electrolyte changes from weakly acidic to alkaline, when the free acid is consumed. Due to the proton discharge at the cathode, one can use an undivided cell; except when easily reducible functional groups, as the nitro group, are present, a divided cell with a diaphragm is needed. As solvents, mostly methanol, but also aqueous methanol or aqueous AN is used. As an anode material, universally platinum in thin foils or as net is applied. Glassy carbon, hard graphite [218], and, as recently shown, the BDD electrode can be additionally used. The latter has been shown to produce the same yield of dimer as the platinum electrode, when the electrolysis of a water-insoluble acid is conducted in water and the electrolyte is emulsified by sonication (see also Chapter 8) [219,220]. For the use of the BBD electrode in phenol oxidations, see Section III.B.2. Soft graphite is not a useful anode material as it promotes further oxidation of the radical to a carbocation, thus disfavoring the radical reaction. As a cathode material, often stainless steel or platinum is used. The former should be applied for substrates with double or triple bonds to avoid an unwanted cathodic hydrogenation. As an electrolysis cell, a beaker-type cell with platinum foils supported on a Teflon frame or a flow-through cell have been applied by the author [2b,221]. Kolbe electrolysis of acetic acid was also performed in the gaseous state using a cell with a permeable electrolyte membrane (PEM) [222]. Problems sometimes arise because of deposits of sparingly soluble products, insoluble polymers, inorganic oxides, or salts being formed at the anode surface; these nonconducting layers can strongly decrease the current and in this way increase the duration of the electrolysis. The addition

© 2016 by Taylor & Francis Group, LLC

749

Oxidative Coupling

of small amounts of a less polar solvent (THF, dioxane, 1,2-dimethoxyethane, pentane), freezing out the product, periodical extraction of the electrolyte, mechanical cleaning of the anode, or short polarity reversals of the anode can be helpful in such cases. External temperatures of −20 to 50°C are regularly applied. At higher temperatures, a conversion of the acid into the electro-inactive methyl ester has to be kept in mind [223]. For more experimental details, the named reviews [216] and Chapter 33 should be consulted. 2. Homocoupling The anodic oxidation of identical carboxylic acids affords homocoupling products. Despite the high discharge potential of the carboxylic acid, a fair number of functional groups are tolerated. These are alkyls, arylalkyls, and groups (CH2)nX with X = COR or CO2R and n > 1 and with X = OAc, NHAc, or Hal and n > 4. The stereochemistry of the products indicates that adsorption of saturated alkyl radicals is unimportant. Carboxylates, which are chiral and nonracemic at the α-position, totally lose their optical activity in heterocoupling [224,225]. This racemization indicates either a free radical as intermediate or its fast adsorption–desorption at the anode. Polar substituents can be handled without protection because nonpolar radicals are involved. This saves steps for protection and deprotection of substrates, whose substituents can react with strong bases, nucleophiles, and electrophiles. Homocoupling via Kolbe electrolysis is a unique and attractive method for the synthesis of symmetrical compounds. A great number of homocoupling reactions have been tabulated in Reference 216. Furthermore, Table 18.5 and the structures in Scheme 18.77 show some selected examples. In general, mainly the substituent in α-position is critical for the yield of the coupling product. Electron-donating substituents (phenyl, vinyl, halo, or amino substituents, and more than one alkyl group) more or less shift the reaction toward products that originate from carbenium ions being formed by further oxidation of the radical (see Section IV.B.6). On the other hand, electron-attracting groups (cyano-, carbonyl substituents or carboxylic acid derivatives) or hydrogen favor dimerization. Carboxylic acids with a double bond in α-position form only small amounts of dimer or no dimers at all. These acids include aromatic acids, pyridine carboxylic acids, and α,β-unsaturated carboxylic acids. Reasons are the slow decarboxylation of the aroyloxy, pyridoyloxy, and acryloyloxy radical and in the latter case additionally the fast radical polymerization of acrylic acid derivatives as competing reaction. TAbLE 18.5 Anodic Homocoupling of Carboxylic Acids Entry 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11.

Carboxylic Acid

yield (%)

References

CH3(CH2)nCO2H, n = 5–15 RO2C(CH2)nCO2H, n = 4–16, R = CH3, C2H5 (CH3)2CH–CH(CO2Et)CO2H AcO(CH2)nCO2H, n = 3–5 F(CH2)nCO2H, n = 4–10 EtCO(CH2)4CO2H (Z)–CH3–(CH2)7CH═CH(CH2)7CO2H (oleic acid) (Z)–CH3–(CH2)7CH═CH(CH2)11CO2H (erucic acid) CF3(CF2)7CO2H CH3O2C(CH2)7CO2H

60–90 45–95 67 73–83 45–70 75 75 70 61 80 60–65

[216b] [216b] [226] [227] [228] [229] [230] [230] [230] [230] [231]

O CH3 P (CH2)nCO2H t-C8H17

© 2016 by Taylor & Francis Group, LLC

n = 1, 2

750

Organic Electrochemistry 40% (MeOH. Pt) 50%

81%

CO2H O

CO2H

CO2H AcO

76 [234]

H

74 [232]

75 [233]

CH2OAc AcO

62% (Pyr, H2O, Et3N)

EtO2C

EtO2C

CO2H

MeO2C

CO2H

AcO 28%

77 [235]

CO2tBu

OAc CH2

78 [236] O

65%

52%

CH2 AcO OAc

O

OAc

CO2H O

93% (SiO2 · pip)

F3C CH3O

CO2H

F

52% (SiO2 · pip)

F

R 15–79% R1

F 83 [217]

CO2H R2

R1: octyl, iBu, iPr R2: H,Me R3: H

86 [241, 242]

Si

CO2H

CO2H

F

F

3

62%

72%

CO2H

OCH3

82 [217]

CH2OAc 81 [238]

80 [230]

79 [237]

84 [239] Boc

CO2H 30% MeO2C 87 [243]

85 [240] O

HN OtBu CO2H

88 [244]

38%

SCHEME 18.77 Anodic homocoupling of different carboxylic acids; the arrows indicate the carbon atom, where the homocoupling occurs; the percentage is the coupling yield.

The dimerization of half-esters of diacids (Table 18.5, entry 2) is also of industrial interest because in this way 1, n-diesters for the preparation of polyesters are easily accessible [245]. Methyl hydrogen azelate, available by double bond cleavage from methyl oleate, has been coupled to the 1,ω-C16 -diester. The diester has been converted into homomuscone (80) (Scheme 18.77), using an acyloin condensation and deoxygenation, and a subsequent 1,4-addition of lithium dimethylcuprate [230]. Hydroxy and amino acids can be dimerized in moderate to good yields, when the substituents are not in α- or β-position and when they are protected against oxidation by acylation (Table 18.5, entry 4; Scheme 18.77: 88) [244]. Efficient syntheses of substituted succinic acids from substituted methyl malonates have been developed in the past [246]; a more recent application is the coupling of 78 (Scheme 18.77) as part of a semibullvalene synthesis [236]. While keto carboxylic acids can be dimerized satisfactorily (Table 18.5, entry 6), the corresponding aldehydes couple poorly. However, good yields can be obtained in these cases, when the acetals, for example 79 (Scheme 18.77), are electrolyzed instead [237].

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751

Oxidative Coupling

1,4-Disilyl compounds, like the dimer of 85 in Scheme 18.77, and others with a variety of substituents at the silicon atom were obtained by Kolbe electrolysis of the corresponding β-silyl carboxylic acids in 60–70% yield. The α-silyl group in the intermediate obviously stabilized the radical and led to good dimer yields by retarding the competing further oxidation. Polymerization could be suppressed by using 1,2-dimethoxyethane as cosolvent. The coelectrolysis of this additive apparently helped to inhibit the polymerization. Kolbe dimers of [2.2.2]-bicyclooctane dicarboxylic acid half-esters 87 (Scheme 18.77) were applied for the synthesis of rigid, rod-shaped carbon skeletons. The homocoupling of carboxylic acids has been successfully used in the synthesis of further target molecules. Kolbe electrolysis of 74 is part of a (+)-α-onocerin synthesis [232], the electrolysis of 75 afforded a dimer with two quaternary C-atoms separated by two methylene groups and was used to study physical properties of such type of compounds [233], and 2,6,10,15,19,23-hexamethyltetracontane, a base of cosmetics and drugs, has been synthesized from 76 [234]. The cyclopropylcarboxylic acid 77 could be coupled to a dicyclopropyl compound [235]. C-Disaccharide 81, which is a potential glucosidase inhibitor, is accessible in few steps [238]. 3-Alkenoic acids 89 dimerize to a mixture of three 1,5-dienes 91a–c (Scheme 18.78); the dimers arise by 1,1′-, 1,3′-, and 3,3′-coupling of the intermediate allyl radical 90. When the 3-position of the allyl radical is increasingly shielded, the portion of 3-coupling decreases. The relative amount of the 1,1′-dimer thus can vary from 52% in 3-dodecenoic acid to 76% in 3-cyclohexylidene–propanoic acid. The configuration of the nonterminal double bond is retained to a high degree (~90%) [242]. With 3-alkenoic acids the dimer yield can be improved by neutralization of the acid with a tertiary amine [247]. (Z)-4-Enoic acids partially isomerize to products with a (E)-configuration. Results from methyl- and deuterium-labeled carboxylic acids support an isomerization by way of a reversible ring closure to a cyclopropylcarbinyl radical. (Z)-n-Enoic acids with n > 5 fully retain their configuration [248]. Using 6-alkenoic acids 92, the intermediate radical 93 partially cyclizes to a cyclopentylmethyl radical 94; the radicals 93 and 94 couple to two homodimers and one heterodimer (Scheme 18.79) [248]. See also Section IV.B.5 for such cyclization. R R

91a

1,1' CO2–

R

–e–, –CO2

89

3

1

1,3'

R R

91b

3,3'

90

R R

SCHEME 18.78

Regioisomeric 1,5-dienes from 3-alkenoic acids.

CH2=CH–(CH2)4CO2–

–e–, –CO2 40%

CH2

CH2=CH–(CH2)3CH2

92

93

CH2=CH–(CH2)8–CH =CH2 42%

SCHEME 18.79

91c

CH2=CH–(CH2)5

(CH2)2

37%

21%

Partial cyclization of radicals from a 6-alkenoic acid.

© 2016 by Taylor & Francis Group, LLC

94

752

Organic Electrochemistry

4 R1CO2H + 4R2CO2H

–e– –CO2

4 R1 + 4 R2

R1–R1 + 2 R1–R2 + R2–R2

SCHEME 18.80 Unsymmetrical and symmetrical products in the anodic decarboxylation of two different carboxylates.

3. Heterocoupling Heterocoupling (cross-coupling) of two different carboxylates allows the synthesis of unsymmetrical compounds (Scheme 18.80). However, as the intermediate radicals combine in most cases statistically, the cross-coupling product contains two symmetrical dimers as major side products. To increase the yield of the wanted heterodimer, the easier available and less costly acid is used in an up to 10-fold excess. This way, only two major products are formed, which makes the isolation of the heterodimer easier. Additionally, more costly acid is incorporated to a large extent into the heterodimer. The two acids should be chosen in a way that allows separating the excess homodimer by distillation or crystallization. Problems due to passivation or due to a follow-up oxidation of the radicals to carbocations are often less pronounced in cross-coupling. Cross-coupling products are also obtained in moderate to good yields with a base supported on a solid polymer. The base could be recycled even when high current densities were applied [249]. Due to its simple procedure and broad scope, cross-coupling by Kolbe electrolysis has found wide application despite the formation of symmetrical dimers as side products. Cross-coupling of carboxylic acids allows the synthesis of rare and new fatty acids, pheromones, chiral building blocks, C-glycosides, or nonproteinogenic amino acids. Selected examples of cross-coupling of carboxylates are given in Table 18.6 and in Schemes 18.81 and 18.82. Entries 2 and 13 in Table 18.6 give examples for the preparation of ω-bromofatty acids. Entries 4 and 5 display chain extensions of unsaturated fatty acids by coelectrolysis with half-esters of diacids. In entries 6 and 7, the preparation of trimethylsilyl compounds with a long alkyl chain by coelectrolysis of trimethylsilyl acetic acid with fatty acids is demonstrated. Entries 8 and 9 show the synthesis of fatty acids that are partially perfluorinated. Entries 10 and 11 give examples for the preparation of long-chain fatty acids with a terminal double bond. Entries 15–17 display the coupling to C-glycosides that are potential enzyme inhibitors. Electrolysis of the ketogulonic acid in entry 15 without a coacid affords no radical dimer but nearly quantitatively the non-Kolbe product [238]. In the presence of a coacid, the non-Kolbe product is still the main product, but a considerable part of the carbohydrate radical can now be trapped prior to oxidation by the radical of the coacid to form the heterocoupling product. In entry 18 (Table 18.6), homologues of dihydro-12-oxophytodienoic acid and jasmonic acid are prepared in one step and a flexible synthesis. Heterocoupling has been used for the extension of the carbon chain in fatty acids [216b]. Furthermore, the method has been applied for the synthesis of pheromones [263]. Some examples are displayed in Scheme 18.81. Muscalure (95), the pheromone of the housefly, has been synthesized by coelectrolysis of heptanoic acid with oleic acid in 14% yield [264]. The moderate yield could be increased to 80% by using a 10-fold excess of heptanoic acid and a lower temperature [265]. Muscalure has been also prepared by coelectrolysis of erucic acid with propionic acid in 59% yield [266]. Furthermore have been prepared: the antagonist of muscalure (Z)-11-heneicosene [265], looplure (96) [267], brevicomin (97) [268], disparlure (98) [269], optically active compound 100, a Trogoderma pheromone [270], and pheromones with a diene or triene structure, such as compound 99 [271]. The chiral acid 101 is one of the building blocks prepared for the total synthesis of the antibiotic myxovirescine [272]. Furthermore, chiral building blocks have been obtained by heterocoupling reactions with (S)-2-(1,1dimethylethyl)-1,3-dioxolane-5-oxo-4-acetic acid 102 [273]. In addition, alkyne carboxylic esters

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753

Oxidative Coupling

TAbLE 18.6 Anodic Cross-Coupling of Two Different Carboxylates R1COOH 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15.

CH3O2C(CH2)4CO2H EtO2C(CH2)8CO2H CH3O2C(CH2)7CO2H CH3O2C(CH2)7CO2H CH3O2C(CH2)7CO2H (CH3)3SiCH2CO2H (CH3)3SiCH2CO2H C6F13(CH2)2CO2H C8F17(CH2)2CO2H CH2═CH(CH2)8CO2H CH2═CH(CH2)8CO2H H3CONH(CH2)10CO2H Br(CH2)10CO2H [EtO]2P(O)CH2CO2H O

O O CO2H

O O 16. O

O

O

CO2H

O O 17.

HO2C

O

O 18.

OMe

R2COOH

yield R1–R2

References

CH3(CH2)nCO2H n = 2, 4, 8, 12 Br(CH2)10CO2H CH3(CH2)16CO2H CH3(CH2)7CH═CH(CH2)7CO2H CH3(CH2)11CH═CH(CH2)7CO2H CH2═CH(CH2)8CO2H CH3(CH2)10CO2H CH3O2C(CH2)7CO2H CH3O2C(CH2)7CO2H CH3O2C(CH2)5CO2H CH3O2C(CH2)7CO2H CH3O2C(CH2)8CO2H CH3O2C(CH2)7CO2H CH3O2C(CH2)8CO2H C11H23CO2H Br(CH2)10CO2H CH2═CH(CH2)9CO2H

12 – 48% 54% 36% 51% 47% 54% 80% 41% 56% 40% 40% 32% 79% 27% 33% 33% 31%

[250] [251] [252] [252] [252] [253] [253] [252] [252] [254] [255] [256] [257] [255] [258] [259] [259]

C7H15CO2H C13H27CO2H (Z)-CH3(CH2)7CH═CH(CH2)7CO2H

70% 51% 46%

[260]

C7H15CO2H

40%

[258]

RO2C(CH2)nCO2H R = Me, Et; n = 4 – 7

10 – 43%

[261]

RCO2H R = C2H5, C5H11, C8H17 C11H23, C15H31

33 – 60%

[262]

O

O CO2H

19.

CO2H N H

with different chain lengths can be obtained [274]; they can be hydrogenated selectively to either (E)- or (Z)-pheromones. Furthermore, useful intermediates for the synthesis of dicarba analogs of cystine peptides have been prepared from l-glutamate [275]. Orthogonally protected 2,5-diaminoadipic acid, 2,6-diaminopimelic acid, and 2,7-diaminosuberic acid derivatives bearing up to four different protecting groups are prepared in one step by mixed Kolbe electrolysis [276]. Branched hydrocarbons for use in emollients and cosmetic compositions have been prepared by anodic crosscoupling of a mixture of linear and branched C6 –C22 fatty acids [277]. Anodic cross-coupling of fatty acids has been recently reviewed [278]. A heterocoupling reaction has been applied in the synthesis of (±)-nephromopsinic acid [279], and for the preparation of 3-alkyl-substituted indoles (Table 18.6, entry 19). Carbohydrates without an electron-donating substituent at the radical center as 103 yield heterocoupling products with different coacids in fair to good yield (Scheme 18.82) [260].

© 2016 by Taylor & Francis Group, LLC

754

Organic Electrochemistry 80% H

68%

H

H3C(CH2)7

H

(CH2)7-(CH2)5CH3

H9C4

O O

(CH2)2-(CH2)4OAc 96

95 48%

33%

H

97 52%

62% O

H H3C(CH2)-(CH2)2

H (CH2)2-(CH2)2CH(CH3)2

(CH2)7-C4H9

98

99 56% tBu

20%

O O

CO2H

2

H

5

2

TBDMSO CO2H 101

OH

100

O

102

SCHEME 18.81 Cross-coupling as a key step in the synthesis of selected pheromones, another biologically active compound, and chiral building blocks; the arrows indicate the bond that is formed in the heterocoupling reaction and the percentage is the yield of the heterocoupling. CO2H H

H

AcO

H

H

OAc

H

OAc

R RCO2H –e–, –CO2

CH2OAc 103

H

H

AcO

H

H

OAc

H

OAc CH2OAc

R C5H11 C7H15 MeO2C(CH2)7

SCHEME 18.82

Yield (%) 69 54 46

Heterocoupling of 3,4,5,6-tetra-O-acetyl-2-deoxy-d-gluconic acid (103).

4. Diastereoselective Coupling Enantioenriched carboxylates with a nonracemic stereogenic center in the α-position totally lose their optical activity in heterocoupling [224,225]. This result indicates that in anodic decarboxylation, either a free radical or a radical that undergoes fast desorption–adsorption at the anode is involved. This is different, if radicals are generated in pairs from diacyl peroxides. If the diacyl peroxide from optically pure ethyl ethylmethylmalonate and dodecanoic acid is photolyzed at −78°C, the cross-coupling product, ethyl 2-ethyl-2-methyl-tridecanoate, is formed in 17% yield and 60% ee with retention of the configuration [280]. a. Facial Selectivity Due to a Chiral Auxiliary Carboxylates with nonracemic chiral auxiliaries have been anodically decarboxylated to explore the face-selective hetero- and homocoupling of the intermediate radicals. 2-Substituted malonamides were subjected to heterocoupling with different coacids; (2R,5R)-2,5-dimethylpyrrolidine served as chiral amido group. Heterocoupling of the acids 104a,b with the coacids 105a–d led to the amides 106a–d with a different diastereomeric excess (Scheme 18.83, Table 18.7) [281].

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755

Oxidative Coupling O

O O

O

N

OH + R

R1

–e–, –CO2

2CO

N

2H

R2

R2

2R-106 a–d

2S-106 a–d

R1

104 a–d

105 a–d

R1

N

+

SCHEME 18.83 Diastereoselective coupling of 2-substituted malonamides 104a–d with the coacids 105a–d to amides 106a–d (R1 and R2 are defined in Table 18.7).

TAbLE 18.7 Diastereoselective Heterocoupling to 106a–d 106 a b c d

R1

R2

yield (%)

de (%)

CH2C6H5 C(CH3)2C2H5 C(CH3)2C2H5 C(CH3)2C2H5

C4H9 C4H9 CH2C(CH3)3 C(CH3)2CO2Et

38 69 42 13

25.9 53.5 69.2 86.1

As side products of 106a–d, compounds arising from hydrogen abstraction and further oxidation of the radical 107a–d originating from 104a–d to a carbocation were found. This was especially pronounced in the case of 106d, where these compounds became the major products. As α-amido radicals assume preferentially the (Z)-conformation [282], the prostereogenic carbon atom in the intermediate radical 107a–d originating from 104a–d is differently shielded by the methyl groups of the chiral auxiliary. For that reason, the coradicals from the coacids 105a–d will approach the radical 107a–d for steric reasons preferentially from the re-face. The diastereoselectivity increases with growing size of R1 and R2, which points to an increasing portion of the (Z)-conformer of 107a–d and a growing steric hindrance for the si-approach (see Scheme 18.84). Besides the chiral auxiliary (2R,5R)-2,5-dimethylpyrrolidine (108), the pyrrolidine 109, the oxazolidine 110, the 2,10-camphersultam 111, and the menthol derivates 112 were used (Scheme 18.85). Thereby moderate yields of the cross-coupling product and a diastereomeric excess of up to 80% de were obtained [283]. b. Facial Selectivity Due to a Stereogenic Carbon Atom in α-Position to the Radical Center Facial selectivity induced by a stereogenic center in α-position to the radical center has been probed with acyclic and cyclic radicals (Scheme 18.86) [284]. The acyclic carboxylate 113 afforded a diastereomeric excess of up to 71.4% (R1 = tBu, R2 = Me; coradical: R3 = C(CH3)2CO2Et) and a 31% yield of the cross-coupling product. Photolysis of peroxides consisting of a dodecanoyl group and a peracetylated tartaric acid or d-gluconic acid group in the solid state at −78°C led to a selective cross-coupling. The chemical yields are 40–60% and the diastereoselectivities 90–95% de [285]. For the acyclic carboxylate 113, a diastereomeric excess (de) of up to 71% (R1 = tBu, R2 = Me, R3 = C(CH3)2CO2Et) and a yield of 31% for the cross-coupling product was found. For the cyclic O R1 N H

R2

107a–d

SCHEME 18.84

Preferred re-approach to the intermediate radical 107a–d.

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756

Organic Electrochemistry CH3O O

N

N

N

108

109

CH3O

110

Ph

R = Ph, tBu, iPr N O2S

O 112

111 R

SCHEME 18.85 Chiral auxiliaries 108–112 for the diastereoselective heterocoupling of carboxylic acids. R3 R2 CO2Me

R1

–e–,–CO2

– 113 COO

OMe

H R2

R3CO2–

H

R1

R2

R2 CO2Me

R1

CO2Me

R1

+

R3 Minor

R3 Major R1

O

CO–2 R3CO2–

R2 R2

O

O

–e–,–CO2

114

+

R2 R2

O

R3

R1

R3

R1

O

trans (major)

R2 R2

O

O

cis(minor)

SCHEME 18.86 Facial selectivity induced by a stereogenic center in α-position to the radical center in an acyclic carboxylate 113 and a cyclic carboxylate 114.

carboxylate 114, a de of up to 88 % (R1 = tBu, R2 = H, R3 = C(CH3)2CO2Et) and a yield of 33% for the cross-coupling product was obtained [248]. 5. Addition of Radicals from Anodic Decarboxylation of Carboxylic Acids to Olefins The anodic decarboxylation of carboxylic acids in the presence of olefins leads to additive dimers 115 and additive monomers 117 (Scheme 18.87, Table 18.8). The products can be rationalized by the following pathway: The radical R• obtained after anodic decarboxylation adds to the alkene to

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757

Oxidative Coupling R– –e–

RCO2– –e– –CO2

Y R

R Y

Y

Y

115

R

117

R

R

Y

R

116

R

–e–

R-R

Y

Y

Y R

+

Nu–

+

R

Nu

R 119

118

120

SCHEME 18.87 Products from the addition of anodically generated radicals to olefins.

TAbLE 18.8 Addition of Radicals from Anodic Decarboxylation of Carboxylic Acids to Olefins Entry

Radical Precursor

Olefin

Conditions

1

EtOOCCOOH

Butadiene

MeOH, Pt

2

MeOOC(CH2)4COOH CH2═CH2

MeOH, Pt

Product (EtOOCCH2CH═CHCH2)2 + isomers MeOOC(CH2)12COOMe

yield References 66%

[286]

95%

[287]

MeOOC(CH2)10COOMe CH2═CHOEt

3

MeOOCCH2COOH

4

CH2═CHPh MeOOC(CH2)4COOH Butadiene

5

MeO2CCH2COOH

6

F3CCOOH

CH2═CHnPr

7

CH3COOH

CH2═C(Me)CHO

8

CF3COOH

CH2═CHOAc

9 10

CH3COOH

CO2H

Dimethylfumarate MeOOCCH2COOH

MeOH, Pt

[MeOOC(CH2)CH(OEt)]2

35%

[288]

MeOH, Pt

[MeOOC(CH2)2CH(Ph)]2

38%

[286] [289]

MeOH, Pt MeCN, NaOH, Pt MeCN, H2O, Pt MeCN, H2O, NaOH MeCN, H2O MeOH, 5% NaOH

(MeOOC(CH2)5CH═CH(CH2)2

47%

(MeOOC(CH2)5CH═CHCH2)5COOMe [CF3CH2CH(nPr)]2

47% 40%

[290]

[CH3CH2C(Me)(CHO)]2

80%

[291]

[CF3CH2CH(OAc)]2

24%

[292]

[MeOOCCH(Me)]2

80% 41%

[293] [294]

53%

[295]

47% 27% 44%

[292]

MeO2C

H

O

O H

11

CO2H

MeOOC(CH2)4COOH MeOH, 5% NaOH

MeO2C(CH2)4 N Ac

N Ac 12

CF3COOH

13

CF3COOH

EtO2CCH═CHCO2Et AN, water, NaOH AN, water, NaOH O O

N Et

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[EtOOCCH(CF3)]2 [EtOOCCH(CF3)CH(COOEt)]2

F3C O

CF3 N Et

O

[292]

758

Organic Electrochemistry

yield the primary adduct 116, which can dimerize to the additive dimer 115 with a regiospecific head-to-head connection of the two olefins. Furthermore, the radical 116 can couple with the radical R• to form the additive monomer 117. If electron-donating substituents Y in the olefin stabilize a carbenium ion, 116 can be oxidized to the cation 118. The cation can react intramolecularly with a nucleophilic site in R to 119 or intermolecularly with a nucleophile in the electrolyte to afford 120. Satisfactory to good yields of adducts have been found for styrenes (Scheme 18.87, Y = phenyl), conjugated dienes (Y = vinyl), enamines (Y = NR2), and enol ethers (Y = alkoxy), particularly if the olefins are not substituted at the carbon atom in β-position to Y (Table 18.8 and Section IV.C.2, Table 18.13). Nonactivated alkenes react less satisfactorily. Radicals generated by anodic oxidation of carboxylates in the presence of olefins form additive dimers 115 (Table 18.8, entries 1, 3–8) and additive monomers (Table 18.8, entries 9–13). In the electrolysis of methyl malonate with vicinal disubstituted styrenes, the adduct yields decrease with increasing size of the substituent in β-position to the phenyl group: H = 42%, Me = 27%, and Et = 11% [296]. The ratio of additive dimer 115 to additive monomer 117 can be influenced to some extent by the current density. In the electrolysis of trifluoroacetate in AN and H2O in the presence of electron-deficient olefins, additive dimers and additive monomers are obtained. The selectivity can be controlled by current density, temperature, and the substitution pattern of the olefin [297]. Trifluormethylation of aromatic compounds with –M substituents, such as benzonitrile, benzaldehyde, acetophenone, and nitrobenzene, has been achieved in moderate yield by the electrolysis of pyridinium trifluoroacetate in AN [298]. Electrolysis of methyl oxalate in methanol with ethylene under pressure yielded 70–90% of the dimethyl esters of succinic, adipic, suberic, and sebacic acids. Increase of the ethylene pressure or decrease of the current density led to an increase in the portion of higher esters in the product mixture [299]. The influence of mechanism and rate on yield and selectivity in the addition of anodically generated radicals to olefins has been calculated and the prediction tested in preparative electrolyses [300]. Good yields can be obtained with nonactivated alkenes, when the addition proceeds intramolecularly. β-Allyloxypropionates and β-allylaminopropionates cyclize in an intramolecular addition and a subsequent cross-coupling of the exocyclic radical to form 3-substituted tetrahydrofurans and pyrrolidines (Table 18.8, entries 10 and 11). This intramolecular addition has been used to synthesize a precursor of prostaglandin PGF2α (Scheme 18.88) [301] and a branched carbohydrate (Scheme 18.89) [302]. CO2H AcO

O

OEt

R

RCO2H –e–, –CO2

OEt

AcO

a R

R: CH3(CH2)7, 54%, a:b = 3.1:1 R: PhMe2SiCH2, 38%, a:b 4.3:1

+

O OEt O

AcO

b

SCHEME 18.88 Synthesis of a precursor of prostaglandin PGF2α via an intramolecular addition and a subsequent cross-coupling reaction.

AcO AcO

O

H

H O

CH3CO2H CO2H

–e–, –CO2 50%

AcO

O

O

AcO

SCHEME 18.89 Synthesis of a branched carbohydrate via an intramolecular addition and a subsequent cross-coupling reaction.

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The extension of the cyclization reaction from tetrahydrofurans and pyrrolidines to carbocycles led to a sharp decrease in the yield of the cyclized product [303]. The reason for that is the slower cyclization rate of 5-hexenyl radicals compared to 5-(3-oxahexenyl) radicals [304], which favors the competing bimolecular coupling to the acyclic product. Changes in the reaction conditions helped to increase the yield of the cyclization product. The geminal dialkyl effect, namely, the introduction of two geminal methyl groups into the 3-position of the 5-hexenyl radical, increases the cyclization rate of the 5-hexenyl radical [304] and leads to a higher portion of cyclization product [303]. The current density controls the concentration of radicals at the electrode surface. A low current density favors the monomolecular cyclization against the bimolecular coupling to an acyclic product. A decrease in the current density by a factor of 30 increased the ratio of cyclic to acyclic product by a factor of 10 [303]. The electrophilicity of the double bond also influences the addition rate of the mostly nucleophilic radical. With vinylic electron-attracting groups, the yield of the carbocyclic compound could be increased to more than 70% (Scheme 18.90, Table 18.9) [303]. This improvement has been extended with further examples including six-membered carbocycles as well as tetrahydrofurans and tetrahydropyrans; good to excellent yields have been obtained by using ethoxycarbonyl- and Boc-substituted double bonds and geminal substitution (Schemes 18.91 and 18.92, Tables 18.10 and 18.11) [305]. A radical tandem cyclization, consisting of two carbocyclizations and a terminating cross-coupling reaction, has been achieved in the coelectrolysis of sodium acetate with carboxylate 121 (Scheme 18.93) R3 CO2H

R1

R3CO2H

R1

–H+, –e–, –CO2 R2

SCHEME 18.90 double bond.

R2

R2

R2

Improved carbocyclization of 5-hexenyl radicals with electron-attracting groups R1 at the

TAbLE 18.9 yields in the Carbocyclization of 5-Hexenyl Radicals with Electron-Attracting groups R1 at the Double bond Entry 1 2 3 4

R1

R2

R3

yield (%)

CN CO2Et CN COCH3

CH3 CH3 H H

(CH2)4CO2CH3 (CH2)4CO2CH3 (CH2)4CO2CH3 CH3

75 76 71 71

R1

R1 CO2H O n

SCHEME 18.91 double bond.

O

R2CO2H

R2 O

–H+, –e–, –CO2 n

O

Carbocyclizations of 5- and 6-alkenyl radicals with electron-attracting groups R1 at the

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Organic Electrochemistry R1

R1 CO2H

n

SCHEME 18.92

R2 O

R2

R3

R3CO2H

R2

–H+, –e–, –CO2

n

O

R2

Cyclizations to substituted tetrahydrofurans and tetrahydropyrans.

TAbLE 18.10 yields in Carbocyclizations of 5- and 6-Alkenyl Radicals with Electron-Attracting groups R1 at the Double bond Entry 1 2 3 4 a

R1

R2

n

yield (%)

Boc Boc CN Boc

CH3 (CH3)2CH CH3 CH3

1 1 1 2

90 75 65 64a

Additionally 16% acyclic product.

TAbLE 18.11 yields in the Cyclizations to Substituted Tetrahydrofurans and Tetrahydropyrans Entry 1 2 3 4

R1

R2

CO2Et CO2Et CO2Et CO2Et

H H H H

R3

n

yield (%)

CH3 CH3O2C(CH2)2 CH2═CH–CH2

1 1 1 1

86 90 89 87

2

89

O CH3 5

CH3

CO2Et

O

O CH2

CH3

1. LDA 2. I(CH2)3CO2Na(85%) 3. C4H7MgBr + OEt 4. H (74%)

O

121 O

O

CO2H O

5 eq. CH3CO2H +

–e–, –CO2

122 42% 2 diastereomers 2.7:1

SCHEME 18.93

+

123

124

15%

8%

Radical tandem cyclization to tricyclic products.

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having two double bonds with 1,5-distances to the intermediate radicals. This reaction provides a short synthetic sequence to tricyclic products, for example, triquinanes [306]. The selectivity for the formation of the tricyclic, bicyclic, and monocyclic products 122–124 could be predicted by using a mathematical simulation based on the proposed mechanism and a reasonable choice of rates. These radical cyclizations can also be achieved by tin-mediated radical chain reactions of alkenyl iodides [304]. However, the electrochemical cyclization has advantages: it avoids the toxic tin hydride and forms two C–C bonds in a one-pot reaction. Thereby, the second C–C bond can be created simply and with a wide variety of groups by the suitable choice of the coacid. 6. Oxidation to Carbocations In anodic decarboxylation of carboxylic acids, products arising from radicals (Kolbe electrolysis) and from carbocations (non-Kolbe electrolysis) are formed. The radical pathway is favored by a high current density, a smooth anode surface (platinum, glassy carbon, BDD electrode), an acidic electrolyte (5–10% neutralization of the acid), and hydrogen or electron-withdrawing substituents in the α-position of the carboxylic acid. The route to carbocations is supported by a low current density, a rough electrode surface (soft graphite), additives (e.g., perchlorate anion, Cu2+), and electron-donating substituents in the α-position. Intermediate radicals with ionization potentials above 8 eV lead preferentially to coupling products, while those with ionization potentials below 8 eV are further oxidized to carbenium ions [307]. The anodic decarboxylation of 3-oxanonanoic acid and 3-oxapentadecanoic acid in methanol leads exclusively to products of the non-Kolbe electrolysis. To investigate whether the outcome of this electrolysis can be shifted toward products of the Kolbe electrolysis, the influence of coelectrolysis, solvent, current density, degree of neutralization, and chain length of the alkoxy group on the anodic decarboxylation of the two acids has been investigated. The results show the Kolbe coupling product can be favored against the non-Kolbe product by an extended alkyl chain in the alkoxy group, coelectrolysis with long-chain fatty acids, ethanol or dimethylformamide as solvent, and a high current density [308]. The non-Kolbe electrolysis has found many synthetic applications. The carboxylic group in an acid can be replaced this way by a hydroxy, alkoxy, or amino group or a double bond. Furthermore, cationic rearrangements or β-cleavages can be induced [216d,g]. The non-Kolbe electrolysis leads to precursors for C–C coupling reactions, but usually does not form directly a new C–C bond via a coupling reaction.

C. ANIONS FROM CH-ACIDS AS RADICAL SOURCE 1. Coupling Besides carboxylates, other organic anions can be coupled at the anode via radicals. Anions of CH-acids couple in satisfactory yields (Table 18.12, entries 1–6). With some substrates, the yield can be substantially improved with sodium iodide as supporting electrolyte or additive. The iodide anion probably acts as mediator. Thereby, dehydrodimers, tetrasubstituted ethylenes, and cyclopropanes can be obtained in up to 90% yield [317]. Electrolysis of methylene malonates in the presence of sodium iodide in an undivided cell results in 50–70% yield of stereoisomeric cyclobutane tetracarboxylates. Thereby at first, the activated methylene compound is reductively dimerized at the cathode and subsequently the two CH-acids undergo a mediated intramolecular coupling [318]. Triacetylmethane forms in a remarkable and stereoselective reaction in high yield a bicyclic trioxabicylooctane (Table 18.12, entry 5). Furthermore, anions of nitroaliphatic compounds are coupled to give vicinal dinitroalkanes (Table 18.12, entry 6). Grignard compounds and borates (Table 18.12, entries 7 and 8) couple to give alkanes; alkyne anions can be coupled to give dialkynes (Table 18.12, entry 9).

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TAbLE 18.12 Anodic Coupling of Carbanions Entry

Precursor for Carbanion

1

(R O2C)2CHR R1 = Me,Et R2 = C1–C8, Ph, allyl, Bzl (EtOOC)3CH

2

1

2

3

n

Conditions

Me

10–90

[309]

H2O, acetone NaOH, EtOH AN, NaOR

[(EtOOC)3C]2

50

[310]

50–98

[309]

69

[311]

92

[312]

n

(RO2C)2C C(CO2R)2 R = Me, Et n = 3,4

O H

N

6 7 8 9

2-Nitrobutane RMgBr R = C5H11, C18H37, Ph R 3B R = C5H11–C8H17 Phenylethyne

O

N

2

Me

Me

(MeCO)3CH

Me N

O

O 5

References

[(R O2C)2C(R )]2 2

pH 1, H2O

Me N

O

yield (%)

AN, NaOR

(RO2C)2CH CH(CO2R)2 R = Me, Et n = 3,4 4

Product 1

Me O

Et3N, AN

MeCO O MeCO O Me O

Me COMe Me

25% NaOH Et2O

3.4-Dimethyl-3,4- dinitrohexane R–R

70 54–60

[313] [314]

MeOH, KOH

R–R

24–82

[315]

THF, LiClO4

1,4-Diphenyl-1,3-butadiyne

35

[316]

The electrochemical oxidation of carbanions, p-Me-C6H4-C(Z1Z2) – with (Z1 = Z2 = CN; Z1 = CN, Z = COOEt; Z1 = Z2 = COOEt), was studied in AN and Bu4NBF4 by voltammetry and macroscale electrolysis. The carbanions, generated by cathodic reduction of the conjugated acids, show a chemically irreversible oxidation peak, whose position reflects the relative basicity of the anions. Kinetic data indicate that neutral radicals are generated at the electrode and that a fast radical–radical coupling is the most effective dimerization process for all the substrates. Homodimers are formed in high yield after exhaustive, one-electron macroscale electrolysis [319]. 2

2. Addition The oxidation of anions from 1,3-dicarbonyl compounds in the presence of olefins leads to different products depending on the structure of the olefin and the applied anode potential. In the presence of butadiene, only the additive dimer 115 (Scheme 18.87) is obtained (Table 18.13, entry 8); on the other hand, with ethyl vinyl ether, only disubstituted monomers 119 and 120 arise (Table 18.13, entries 1–3). With styrene, 119 and 120 are found as main products and 115 as side product (Table 18.13, entries 4, 7). These results indicate that at a potential between 0.6 and 1.4 V, the intermediate ethoxymethyl radical 116 (Y = OEt) is rapidly oxidized to 118, while the oxidation of the allyl radical 116 (Y = vinyl) is much slower. If styrene and oxygen are present, the electrogenerated radical adds successively to both (Table 18.13, entry 5). With Mn(OAc)3, generated by oxidation of Mn(OAc)2 as mediator, a tandem reaction consisting of an intermolecular radical addition followed by an intramolecular electrophilic aromatic

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Oxidative Coupling

TAbLE 18.13 Anodic Addition of Carbanions to Olefins Entry 1

Radical precursor MeOCH2COMe

Olefin

Conditions

CH2═CHOEt

Product

MeONa, MeOH, Pt

O

Me 2

3

(MeOOC)2CH2

CH2═CHOEt

MeONa, MeOH, Pt

CH2═CHOEt

O

EtOH, EtONa

yield

References

36%

[320]

37%

[320]

79%

[321]

85%

[322]

79%

[322]

67%

[108]

40%

[320]

46%

[320]

80–86%

[323]

40–91%

[324]

37–59%

[325]

OEt

O

MeOOC

OEt OMe

MeOOC O OEt O OEt

O O

OEt O 4

CH2═CHPh

O

MeCN, Et4NOTos

O Ph O

O 5

O

CH2═CHPh, O2

MeCN, Et4NOTos

O Ph O

O

O 6

CH3COCH2CO2Me

N

MeOH, O MeONa

N

MeOC 7

(MeO2C)2CH2

8

(MeO2C)2CH2

9

CH3COCH2CO2Et

10

PhCH2CH(CO2Me)2

11

PhCOCH2NO2

CH2═CHPh

Butadiene

MeOH, MeONa, Pt

OH O

CO2Me + isomer

MeO2C MeO MeO

O

Ph

MeOH, [(MeO2C)2CHCH2CH═CHCH2]2 MeONa, + isomer Pt HOAc, CO2Et R2 R2 EtOAc, 1 R NaOAc Me R1 O R1 = Propyl, (0.1 eq.), Pentyl, t-Bu Mn(OAc)2, R2 = H, Me C-anode KOAc, R CO2Et H R = Hexyl HOAc CO2Et TMS, Ph 0.25 eq. Mn(OAc)2, R C-anode HOAc, CH2═CHR O R = Bu, CH2OAc Bu4NBF4 R Ph Mn(OAc)2, C anode N+ O 60oC O–

(Continued)

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Organic Electrochemistry

TAbLE 18.13 (Continued) Anodic Addition of Carbanions to Olefins Entry

Radical precursor

Olefin

Conditions

12

(CH3)2CHNO2

CH2═CHPh

MeOH, MeONa, Pt

Product

yield

13

nBuMgBr

CH2═CHPh

14

(C6H13)3B(OMe)•

Butadiene

Et2O, [nBu-CH2-CH(C6H5)]2 LiClO4, Pt MeOH, (C6H13CH2CH═CHCH2)2 MeONa, C C6H13CH2CH═CHCH2C6H13 C6H13CH2CH═CHCH2OMe + isomers

NO2 Ph

References

43%

[326]

29%

[327]

5% 14%

[328]

OMe

substitution can be accomplished (Scheme 18.94) [329]. Further, Mn(III)-mediated additions of 1,3-dicarbonyl compounds are shown in Table 18.13 entries 9–11. Mediated by in situ generated Mn(III), methyl dibromoacetate, trichlorobromomethane, perfluoroctyl iodide, dibromomalonate, and active methylene compounds have been added to olefins via radicals [330]. Recently, a large number of anodic additions of CH-acids to olefins mediated with Mn(III) or Ce(IV) have been compiled [317]. Sorbic acid precursors have been obtained in larger scale and high current efficiency by a Mn(III)/ Cu(II)-mediated oxidation of acetic acid and acetic anhydride in the presence of butadiene [331]. Also other anions, such as the 2-nitropropanate anion, Grignard reagents, and borates, can be added to olefins (Table 18.13, entries 12–14). In compounds bearing a silicon and a tin atom on the same carbon atom, the tin–carbon bond can be selectively cleaved at the anode and the resulting intermediate added to a silyl enol ether or an allylsilane (Scheme 18.95) [332]. CO2Et R3

R3

R2

R1

+ X

CO2Et

a –e–

X

R2

R1

R1 = H, F, Me; R2 = H, Me; R3 = H, Me; X = CN, CO2Et a: C anode; NaOAc, HOAc, AcOEt (13:3), 0.2 eq. Mn(OAc)2, 60°C

SCHEME 18.94

Subsequent inter- and intramolecular addition of an anodically generated radical. –e–, DCM, 0.2 M Bu4NBF4

OMe Bu3Sn

SiMe3

+

OMe

SiMe3 99% OSiMe3

SPh + Bu3Sn

O

SiMe3 SPh

–e–, DCM, 0.2 M Bu4NBF4

SiMe3

SiMe3 70%

SCHEME 18.95 Selective cleavage of a tin–carbon bond at the anode and addition of the electrophilic intermediate to double bonds.

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765

V. OUTLOOK C–C bond formation at the electrode consists of two reactions: (a) Electrons are transferred between substrate and electrode forming radical cations or radicals as reactive intermediates and (b) chemical reaction of these intermediates lead to C–C bond formation. Reaction (a) depends on the facility for the transfer of an electron to or from the substrate. This is determined by the energy of the π- or σ-bonds, the lone electron pairs, the interaction of substrate and electrode, and the electrode potential. Reaction (b) is controlled by the reactivity of the intermediate and the composition of the reaction layer. Most reactive species being involved in electrosynthesis react in a reaction layer close to the electrode surface. The faster they react, the less they can diffuse away from the electrode and the more the electrode and the reaction layer close to the electrode surface influence the chemical reaction. This can be due to the electrostatic field of the electrode, a further electron transfer to or from the electrode, and a different electrolyte composition and substrate concentration in the reaction layer than in the bulk solution. Much information on the rate and facility of electron transfer at the electrode is available or accessible by experimental methods and quantum chemical calculations (see Chapters 1 to 6). Much less is known on the composition of the reaction layer. Knowledge is emerging from well-supported interpretations on the selectivity in the intramolecular reaction of radical cations with double bonds or in the intermolecular cross-coupling of phenols. (See Sections III.E.1 and III.E.2.) It is to be expected that these analyses will stimulate others in the interpretation of their results and this way widen our knowledge on the reaction layer. Possibly established tools available from bulk reactions as the linear free energy relationship, salt effects, or competition experiments will be applied to further characterize the properties of the reaction layer. Increasing the hydrophobicity of the supporting electrolyte by exchanging LiClO4 for Et4NOTos has led to a lower content of methanol in the reaction layer, and this way has favored an intramolecular cyclization against the methanolysis of the intermediate radical cation [164]. Helpful would be to find, for larger-scale conversions (>0.1 F ), reliable ways to regenerate or retain the activity of the electrode surface or to concentrate catalysts in the reaction layer by using electrostatic attraction or by linking them with flexible tethers to the electrode surface. The principle of redox umpolung to save reaction steps, couple substrates of equal reactivity, and use the same building block as anion, radical, radical ion, or cation should be more systematically investigated. Sometimes one can learn from a related conversion that works well with a chemical oxidizing or reducing agent about the deficits of the electrochemical reaction, or one can use such a species as a mediator of an indirect electrolysis. Worthwhile would be a look on combinations of electron transfer with organocatalysis or transition metal catalysis and also include such combinations into indirect electrolyses. Transition metal catalyst can be oxidized to higher oxidation states, where they can be more powerful oxidants and this way may lead to yet unexplored catalytic reactions. Whenever possible, simple and inexpensive equipment should be used: many anodic oxidations can be done in an undivided cell that simplifies the cells and keeps the cell voltage low and the acidity of the electrolyte constant. A CCE with a cheap direct current (dc) power supply saves a coulometer and an expensive potentiostat. The electrodes should be used as solid plates or in case of thin foils be mounted on Teflon frames. Each glassblower can manufacture a beaker-type cell, and the machine shop can produce a Teflon stopper with holes for the current feeders and, if needed for a Luggin capillary, thermometer and inert gas inlet. The experimental part in a publication should be written up in a way that somebody without expertise in electrochemistry understands the procedure, all that helps to repeat, to evaluate, and possibly to use an electrochemical reaction quickly in most chemical laboratories. At teaching institutions, the basis of electrosynthesis should be presented in lectures and laboratory courses to decrease the barrier against using electrosynthesis.

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Finally, there are a number of advantages—being partly unique—that are offered by electrosynthesis. These are redox umpolung, the use of the same synthetic building block in different reactivity, access to reactions being not available in nonelectrochemical reactions, accord with many rules of green chemistry, high atom economy, cheap reagents, easy scale-up, or facile application in microreactors. However, one should keep in mind that electrosynthesis cannot create miracles, but it can offer alternatives to nonelectrochemical syntheses that may be a better solution for a synthetic problem.

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255. Plate, M.; Overs, M.; Schäfer, H.J. Synthesis 1998, 1255–1258. 256. (a) Gockeln, M. University of Münster, Münster, Germany, 1991; (b) Schäfer, H.J. Eur. J. Lipid Sci. Technol. 2012, 114, 2–9. 257. Jensen-Korte, U.; Schäfer, H.J. Liebigs Ann. Chem. 1982, 1532–1542. 258. Weiper, A.; Schäfer, H.J. Angew. Chem. 1990, 102, 228–230; Angew. Chem. Int. Ed. Eng. 1990, 29, 195–197. 259. (a) Weiper, A., PhD thesis, University of Münster, Münster, Germany, 1990; (b) Schäfer, H.J. Eur. J. Lipid Sci. Technol. 2012, 114, 2–9. 260. Harenbrock, M.; Matzeit, A.; Schäfer, H.J. Liebigs Ann. 1996, 55–62. 261. Schierle, K.; Hopke, J.; Niedt, M.-L.; Boland, W.; Steckhan, E. Tetrahedron Lett. 1996, 37, 8715–8718. 262. Yadav, A.K.; Singh, A.; Prakash, L. Indian J. Chem. Sec. B. 1998, 37B, 1274–1278. 263. Schäfer, H.J. Chem. Phys. Lipids 1979, 24, 321–333. 264. Gribble, G.W.; Sanstead, J.K.; Sullivan, J.W. J. Chem. Soc. Chem. Commun. 1973, 735–736. 265. (a) Seidel, W. PhD thesis, University of Münster, Münster, Germany, 1980; (b) Schäfer, H.J. Eur. J. Lipid Sci. Technol. 2012, 114, 2–9. 266. Yadav, A.K.; Tissot, P. Helv. Chim. Acta 1984, 67, 1698–1701. 267. Seidel, W.; Knolle, J.; Schäfer, H.J. Chem. Ber. 1977, 110, 3544. 268. Knolle, J.; Schäfer, H.J. Angew. Chem. 1975, 87, 777–777; Angew. Chem. Int. Ed. Eng. 1975, 14, 758–758. 269. Klünenberg, H.; Schäfer, H.J. Angew. Chem. 1978, 90, 48–49; Angew. Chem. Int. Ed. Eng. 1978, 17, 47–48. 270. Jensen, U.; Schäfer, H.J. Chem. Ber. 1981, 114, 292–297. 271. (a) Bestmann, H.; Roth, K.; Michaelis, K.; Vostrowsky, O.; Schäfer, H.J. Liebigs Ann. Chem. 1987, 417–422; (b) Steinbauer, M.J.; Ostrand, F.; Bellas, T.E.; Nilsson, A.; Andersson, F.; Hedenstrom, E.; Lacey, M.J.; Schiestl, F.P. Chemoecology 2004, 14, 217–223. 272. Seebach, D.; Maestro, M.A.; Sefkow, M.; Neidlein, A.; Sternfeld, F.; Adam, G.; Sommerfeld, T. Helv. Chim. Acta 1991, 74, 2112–2118. 273. Seebach, D.; Renaud, P. Helv. Chim. Acta 1985, 68, 2342–2349. 274. Seidel, W.; Schäfer, H.J. Chem. Ber. 1980, 113, 3898–3903. 275. Nutt, H.F.; Strachan, H.G.; Veber, D.F.; Holly, F.W. J. Org. Chem. 1980, 45, 3078–3080. 276. Hiebl, J.; Kollmann, H.; Rovenszky, F.; Winkler, K. Bioorg. Med. Chem. Lett. 1997, 7, 2963–2966. 277. Dierker, M. PCT International Application 2006, WO 2006094642 A1 20060914. 278. Schäfer, H.J. Eur. J. Lipid Sci. Technol. 2012, 114, 2–9. 279. Brecht-Forster, A.; Fitremann, J.; Renaud, P. Helv. Chim. Acta 2002, 85, 3965–3974. 280. Feldhues, M.; Schäfer, H.J. Tetrahedron 1985, 41, 4213–4235. 281. (a) Klotz-Berendes, B.; Schäfer, H.J. Angew. Chem. 1995, 107, 218–220; Angew. Chem. Int. Ed. Eng. 1995, 34, 189–191; (b) Klotz-Berendes, PhD thesis, University of Münster, Münster, Germany, 1994. 282. Porter, N.A.; Giese, B.; Curran, D.P. Acc. Chem. Res. 1991, 24, 296–304. 283. (a) Klotz-Berendes, PhD thesis, University of Münster, Münster, Germany, 1994; (b) Letzel, M. PhD thesis, University of Münster, Münster, Germany, 1997. 284. Hauck, M. PhD thesis, University of Münster, Münster, Germany, 1996. 285. Lomölder, R.; Schäfer, H.J. Angew. Chem. 1987, 99, 1282–1283; Angew. Chem. Int. Ed. Eng. 1987, 26, 1253–1254. 286. Schäfer, H.; Pistorius, R. Angew. Chem. 1972, 84, 893–894; Angew. Chem. Int. Ed. Eng. 1972, 11, 841–842. 287. Vasilev, Y.B.; Kanevskii, L.S.; Karapetyan, K.G.; Kovsman, E.P.; Skundin, A.M.; Tarkhanov, G.A.; Freidlin, G.N. Elektrokhimya 1978, 14, 770; Chem. Abstr. 1978, 89, 119649. 288. Schäfer, H.; Stork, L. University of Münster, Münster, Germany, 1977. 289. Fioshin, M.Y.; Salmin, L.A.; Mirkind, L.A.; Kornienko, A.G. Zh. Obshch. Khim. 1965, 10, 594; Chem. Abstr. 1966, 64, 1949d. 290. Brookes, C.J.; Coe, P.L.; Owen, D.M.; Pedler, A.E.; Tatlow, J.C. J. Chem. Soc. Chem. Commun. 1974, 323–324. 291. Chkir, M.; Lelandais, D. J. Chem. Soc. Sec. D, Chem. Commun. 1971, 1369–1370. 292. Renaud, R.N.; Champagne, P.J.; Savard, M. Can. J. Chem. 1979, 57, 2617–2620. 293. Champagne, P.J.; Renaud, R.N. Can. J. Chem. 1980, 58, 1101–1105. 294. Huhtasaari, M.; Schäfer, H.J.; Becking, L. Angew. Chem. 1984, 96, 995–996; Angew. Chem. Int. Ed. Eng. 1984, 23, 980–981. 295. Feldhues, L.; Schäfer, H.J. Tetrahedron Lett. 1988, 29, 2797.

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Oxidative Coupling 296. 297. 298. 299. 300. 301. 302. 303. 304. 305. 306. 307. 308. 309. 310. 311. 312. 313. 314. 315. 316. 317. 318. 319. 320. 321. 322. 323. 324. 325. 326. 327. 328. 329. 330. 331. 332.

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19

Oxidative Substitution and Addition Reactions Ole Hammerich and James H.P. Utley

CONTENTS I. Introduction ............................................................................................................................ 776 II. Formation of Carbon-to-Carbon Bonds ................................................................................. 777 A. Aromatic Hydrocarbons and Alkenes as Nucleophiles .................................................. 777 B. Trifluoromethylation ....................................................................................................... 778 C. Cyanation ........................................................................................................................ 778 D. Carbomethoxylation ........................................................................................................ 780 E. Allyltrimethylsilane as Nucleophile ............................................................................... 781 III. Formation of Carbon-to-Nitrogen Bonds ............................................................................... 781 A. Acetamidation ................................................................................................................. 781 B. N-Cyanomethylation and Amination.............................................................................. 784 C. Nitration .......................................................................................................................... 784 D. Cyanate Ion as Nucleophile ............................................................................................ 785 E. Pyridination .................................................................................................................... 785 F. N-Heterocyclic Compounds as Nucleophiles ................................................................. 786 IV. Formation of Carbon-to-Oxygen Bonds ................................................................................. 787 A. Hydroxylation.................................................................................................................. 787 B. Alkoxylation (Methoxylation) ......................................................................................... 789 C. Acyloxylation (Acetoxylation) ........................................................................................ 795 D. Nitrate Ion as Nucleophile .............................................................................................. 798 V. Formation of Carbon-to-Sulfur Bonds ................................................................................... 798 A. Thiols as Nucleophiles .................................................................................................... 798 B. Thiocyanation ................................................................................................................. 798 VI. Formation of Carbon-to-Halogen Bonds ................................................................................ 798 A. Fluorination..................................................................................................................... 798 B. Chlorination and Bromination ........................................................................................ 799 C. Iodination ........................................................................................................................ 799 References ......................................................................................................................................800 This chapter is dedicated to the late professor Lennart Eberson, who was not only one of our colleagues but also a close friend. Lennart Eberson was a world-leading physical-organic chemist who made seminal contributions to many areas of the subject. These included pioneering achievements in the application and theory of homogeneous and heterogeneous electron transfer reactions of organic compounds. He was one of a small band of scientists who shaped modern organic electrochemistry and who also contributed to this book from the very beginning.

775

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I. INTRODUCTION Anodic substitution and addition reactions are among the most important reactions in organic electrochemistry, and new, interesting examples are added to the literature every year. The overall reactions may be represented as in the following: R E + Nu– C C

(19.1)

R Nu + E+ + 2e–

+ 2 Nu–

(19.2)

C Nu + 2e–

Nu C

The nucleophile [Nu– (or NuH)] may include H2O, ROH, OH–, RO –, RCOO –, NO3–, SCN–, SeCN–, CH3CN, NO2–, N3–, OCN–, pyridine, halide ion, and CN–; the electrophile [E+] in substitution reactions is most commonly H+ but might be a carbocation or an alkyloxonium ion. The cases in which the electrophile is CO2 (the Kolbe reaction and similar) are treated in Chapter 33 and will not be dealt with here. Anodic substitution and addition reactions therefore allow for a net reaction with a nucleophile that is not easily achieved by ordinary chemical means. Specifically, it makes possible the direct formation of a C–O or C–N bond, which is of considerable synthetic interest. The mechanisms of anodic substitution and addition reactions can be divided into direct and indirect processes. In the direct process, oxidation of the organic substrate takes place at the anode with formation of a radical cation as the first step, Equation 19.3, which is followed by reaction with a nucleophile and/or base present in the electrolyte, Equation 19.4 and/or Equation 19.5. The resulting neutral radical, I or II, is usually more easily oxidized than the starting material and undergoes further oxidation at the applied potential to a cation, III or IV, Equation 19.6 or Equation 19.7, which in turn is converted to the final product, V, in a fourth step, Equation 19.8 or Equation 19.9: +

R H R H

+

(19.3)

+ e–

R H

H

H + Nu– (or NuH)

(or

R Nu I

+

R H

+ B

R II

+ BH+

R

) NuH+ I'

(19.4)

(19.5)

Electrode – e–

H

H

R

+R Nu I

Nu III

Solution +R

+ H , –R

(19.6)

H

Electrode –e– R+

R II

IV

Solution

+R

+ H ,

–R

(19.7)

H

H +R

+ B Nu III

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R Nu + BH+ V

(19.8)

777

Oxidative Substitution and Addition Reactions R+ + Nu– IV Nu C C + + Nu–

R

(19.9)

Nu V

Nu C

(19.10)

C Nu

The mechanism composed by Equations 19.3 through 19.6 and 19.8 is observed typically for nuclear substitution in aromatic compounds, whereas side-chain substitutions, for instance, typically proceed according to Equations 19.3, 19.5, 19.7, and 19.9. (For a more thorough discussion of the reactions between radical cations and nucleophiles, the reader is referred to the rich literature on this subject [1].) Essentially, the same mechanistic pattern is found in oxidative substitution driven by high-valent inorganic ions [2–5]. Like in conventional organic chemistry, the competition between substitution and addition reflects whether the cation, in electrochemistry resulting from a −2e–, −H+ process, deprotonates, Equation 19.8, or reacts with a second nucleophile, Equation 19.10. In an indirect process, the anode merely serves as a convenient source of an oxidant, for example, Cl2 by oxidation of Cl– or NO3• by oxidation of NO3–, or of an organic radical, for example, methoxy radical from methoxide ion, which then attacks the organic substrate in a reaction similar to that taking place when the reagent is generated in any other way. The indirect processes also encompass reactions in which an organic compound is oxidized by an electrochemically generated oxidant, a so-called mediator, the reduced form of which is continuously reoxidized and therefore, in principle, needs only to be present in catalytic amounts. In this chapter, we shall emphasize reactions proceeding according to the direct mechanism, since in many cases the indirect processes do not differ in any important respects from analogous homogeneous processes. The subject has been extensively reviewed [6–10].

II. FORMATION OF CARbON-TO-CARbON bONDS A. AROMATIC HYDROCARBONS AND ALKENES AS NUCLEOPHILES Reaction conditions for anodic substitution may often be arranged so that the starting material is the most nucleophilic species present. Consequently, many synthetically useful inter- and intramolecular reactions with carbon-based nucleophiles may be accomplished (see Chapter 18 for details of anodic coupling reactions). An obvious case is the oxidation of electron-rich hydrocarbons in nonnucleophilic media, typically CH2Cl2/R4NBF4 or CH2Cl2 mixed with strong acids such as CF3COOH, CH3SO3H, or CF3SO3H [11,12]. The radical cations formed by one-electron oxidation either dimerize or attack the starting material; an example is given in Equation 19.11. Such couplings have also been achieved using a large-scale capillary gap cell with a graphite anode [12]. CH3 CH3

H3C

CH3

+CH

CH3

2

CH3 –2e–, –H+

H3C CH3

CH3 H3C

CH3

CH2

–H+

H3C

CH3

H3C

(19.11)

CH3

CH3 H3C CH3

Intramolecular coupling of aromatic systems based on this principle has been particularly fruitful. Phenols or aromatic ethers are usually involved: if the activation of the aromatic system is sufficient, the reactions do not require especially nonnucleophilic solvents. In the simplest case, anisole

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Organic Electrochemistry

may be anodically coupled in CH2Cl2/CF3COOH to give, after work-up in the presence of Zn dust, 4,4′-dimethoxybiphenyl in 60% yield [13]. Related to these coupling reactions is the reaction between an anodically generated radical cation, for example, derived from anthracene or 9-phenylanthracene, and a carbon nucleophile such as anisole, toluene, or benzene [14]: CH3 CH3

B.

(19.12)

~100%

–2e–, –2H+

TRIFLUOROMETHYLATION

There is a growing interest in the synthesis of fluorinated organic compounds (see, e.g., Chapter 20), and new methods of preparing, for example, trifluoromethyl-substituted compounds are desired. Oxidation of trifluoromethanesulfonate ion in MeCN or N,N-dimethylformamide (DMF) has been shown to lead to the formation of the trifluoromethyl radical that in a subsequent step may react with a suitable olefin such as 1-phenylcyclohexene, Equation 19.13 [15]. The total yield of CF3 products is 35–70%.

CF3SO2–

–e–, –SO2

CF3

F3C

F3C

C.

(19.13)

+

CYANATION

Anodic cyanation is a reaction of considerable preparative interest since it allows for the direct introduction of a cyano group into an aromatic ring [16–24]. Contrary to the situation in acyloxylation (see Section IV.C) and alkoxylation (see Section IV.B), the product is more resistant toward oxidation than the starting material owing to the strong electron-withdrawing effect of the cyano group, and in principle, good yields should be obtainable. The reaction is often run in MeOH/NaCN [16–20], which gives concurrent methoxylation [25]. This is avoided if MeCN/Et4NCN is used [21]. An important advance is the use of emulsions with phase-transfer agents [23]. The reaction was initially believed to be a homolytic substitution reaction by anodically generated cyano radicals [16]. However, while significant oxidation of CN- takes place at a potential as low as 0.5 V versus SCE, the cyanation process will occur only in the region around or above E1/2 of the substrate. This strongly implies a direct mechanism [17,20,21]. Isomer distribution for cyanation of aromatic compounds has been determined in a number of cases [18–21], and the regioselectivity is often high. This has been put to good use for the preparation of 2-cyanopyrroles and 2-cyanoindoles [24,26,27]: CH3OH, CN– N CH3

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–2e–, –H+

N CH3

CN

64%

(19.14)

779

Oxidative Substitution and Addition Reactions

If the 2-position is blocked, cyanation occurs at the 3-position. Similarly, cyanation of 1-methylpyrazole gives a mixture of the 4- and 5-carbonitrile [28]. The efficient substitution observed for many pyrroles, imidazoles [29], and indoles contrasts markedly the addition reactions observed for furans and thiophenes [30]. However, this difference may only be apparent. At least for two of the cases cited earlier [24], it has been demonstrated [31] that the cyanosubstituted pyrroles arise as a result of elimination during work-up of the 2,5-addition product originally formed. For benzo[b]thiophenes, cyanation leads predominantly to substitution products [32]: CN OH, CN–

CH3 S

CH3

CH3

–2e–, –H+

75%

(19.15)

S

The direct nature of attack of CN– on the radical cations of aromatic compounds has been demonstrated by cyclic voltammetry [22]. The reversible one-electron oxidation of anthracene becomes chemically irreversible in the presence of CN–, and 2F electrolysis gives a mixture of cyano and isocyano addition across the 9,10-position. Interestingly, it appears that cyanation of 9,10-diphenylanthracene gives the 9,10-diphenyl-9,10-dicyano-9,10-dihydroanthracene only [33]. The best conditions for nuclear cyanation involve the use of emulsions, typically CH2Cl2/H2O/ NaCN, with a phase-transfer agent. This has been described in detail for the monocyanation of naphthalene and anisole with isolated yields in the region 50–70% [23]. For dimethoxybenzenes, the yields are somewhat lower [23]. For trans-stilbene and trans-4,4'-dimethoxystilbene, cyanation in an emulsion system gives for the former a mixture of addition and substitution products (including substitution of the vinylic protons), whereas for the latter substitution of the vinylic protons is the predominant reaction [23]. Alkylaromatic compounds do not give any side-chain substitution product under conditions that produce nuclear cyanation products [18–20] (see previous text). Anodic cyanation of tertbutylanisoles leads to mixtures of addition and substitution products [34]. In most cases, cyanation involves loss of a proton, but in some cases, an alkoxy group may be displaced with very good yields of substitution products [21,23]. Thus, 1,2- and 1,4-dimethoxybenzene undergo cyanation by displacement of one of the methoxy groups, presumably via an eCe-type mechanism. On the other hand, anisole and 1,3-dimethoxybenzene do not undergo methoxy displacement but, instead, give nuclear substitution in the usual manner. The presence of an activating group in the ortho and para positions thus seems essential for methoxy group displacement. However, the selectivity is less pronounced when the reactions are carried out in aqueous micellar solution [35,36]. The fate of the displaced alkoxy group is not always clear; the observation of aldehyde and alcohol among the products of anodic cyanation of 4,4′-dialkoxybiphenyls [23] and the proven loss of benzyloxy radical from the radical cation of dibenzyl ether [37] support the mechanism shown in Equation 19.16. The cyanation of the 4,4′-dialkoxybiphenyls, Equation 19.16, shows the use of emulsion systems to good advantage; the products include an important class of liquid crystals. OCH2R

NC

CN– –e–

OCH2R

© 2016 by Taylor & Francis Group, LLC

CN

OCH2R

+ ½RCH2OH + ½RCHO –RCH2O

OCH2R

OCH2R

(19.16)

780

Organic Electrochemistry

Tertiary (cyclic) amines are particularly useful substrates for cyanation [21,23,38–44] with substitution taking place at the carbon atom α to the nitrogen atom as in N-methylpiperidine, Equation 19.17 [38]. Related to these is the α-cyanation of N-(methoxycarbonyl)piperidine [45]. Current–potential curves have been used to show that discharge of, for example, triethylamine occurs prior to the oxidation of the MeOH/NaCN electrolyte. The reaction works well even for complex amines such as benzyl germatranes, Equation 19.18 [46]: CN CN– –2e–, –H+

N CH3

N CH2CN +

(19.17)

N CH3

38%

62% CN N

N O O Ge O

Br

CN– –2e–, –H+

O Ge

O O

(19.18)

Br

For aniline derivatives, substitution often occurs preferentially at a primary carbon atom as in, for example, N-ethyl-N-methylaniline [21]: N

CH3 CH2CH3

CN– –2e–, –H+

N

CH2CN

N

+

CHCH3 36% CN

CH2CH3 64%

D.

CH3

(19.19)

CARBOMETHOXYLATION

An unusual and interesting carbomethoxylation reaction occurs during electrolysis of an arylolefin in MeOH/NaOMe under CO pressure between two Pt electrodes [47]: R

R C CH2

+ CO

+

CH3O–

Pt anode –2e

C CHCOOCH3

–, –H+

(19.20)

It has been shown that the Pt cathode dissolves to give a Pt(CO) complex. This is transformed at the anode to a carbomethoxy–Pt(CO) complex that attacks the arylolefin to give the carbomethoxylation product. Carbomethoxylation of styrene derivatives may also be effectively accomplished by oxidation in AcOH/Ac2O/NaOAc mixtures containing small amounts of Cu(OAc)2 and Mn(OAc)2, the latter serving as a mediator. However, under these conditions, the reaction takes a slightly different course and leads predominantly to cyclized products as illustrated by one, Equation 19.21, of the several examples reported [48]. Related to this reaction is the oxidation of 5-arylpent-1-enes in the presence of, for example, diethyl malonate, Equation 19.22, [48]: H CH2

CH3COOH, (CH3CO)2O, CH3COO– Mn(OCOCH3)2, Cu(OCOCH3)2 75%

O O

© 2016 by Taylor & Francis Group, LLC

(19.21)

781

Oxidative Substitution and Addition Reactions R3

R2

R1

X CH2 COOCH2CH3

+ R3 R1

R2

CH X

CH3COOH, CH3COOCH2CH3, CH3COO– Mn(OCOCH3)2

COOCH2CH3

(19.22)

X = CN, COOCH2CH3

E. ALLYLTRIMETHYLSILANE AS NUCLEOPHILE Allyltrimethylsilane has been used as the nucleophile for the introduction of the α-allyl group in N-(methoxycarbonyl)pyrrolidine, Equation 19.23, and related compounds using a microflow reactor or by taking advantage of acoustic emulsification [49]. Yields are typically in the range 40–70%. In a series of similar substitution reactions, a phenylthio substituent has been replaced by an allyl substituent [50], mostly likely via a carbocation intermediate (see also Equation 19.32). Other examples of the application of allyltrimethylsilane as a nucleophile are based on the so-called cation pool method (see Chapters 9 and 35) that includes the electrochemical generation and accumulation of highly reactive cations, such as the (methoxyphenyl)methylium cation, at low temperature and subsequent reaction with, for example, allyltrimethylsilane to afford (1-methoxybut-3-enyl)benzene in up to 90% yield [51]. A spectacular example of the application of allyltrimethylsilane as a nucleophile is the oxidation of cyclooctatetraene. As seen in the following, the original eight-membered ring is preserved in resulting product, VI [52]: N COOCH3

+

Si(CH3)3 –2e–, –H+

N COOCH3

Si(CH3)3

(19.23) VI

III. FORMATION OF CARbON-TO-NITROgEN bONDS A. ACETAMIDATION Oxidation of anthracene in MeCN containing (CF3CO)2O (to remove water and suppress hydroxylation) leads to the formation of the acetamidation product in 85% yield [53]. Results obtained by cyclic voltammetry supported the mechanism shown in the following: +

CH3CN

–e–

H

+ N C CH3

+ N C CH3

NHCOCH3 H2O

–e–, –H+

© 2016 by Taylor & Francis Group, LLC

–H+

(19.24)

782

Organic Electrochemistry

The rapid reaction of the anthracene radical cation with MeCN is nicely illustrated by the observation of an irreversible, two-electron oxidation peak in the voltammogram at +1.2 V versus SCE, whereas in CH2Cl2/(CF3CO)2O, the oxidation peak corresponds to a reversible oneelectron process. Methyl-substituted benzenes are particularly good substrates for side-chain acetamidation and have featured in numerous product and mechanistic studies [7–10,54–59]; ethylbenzene and isopropylbenzene give other products predominantly [60]. Hexamethylbenzene has been a favored substrate for mechanistic investigations of acetamidation [57,61], and in this case, there is evidence [62] that on time scale of conventional cyclic voltammetry, proton loss is rapid from hexamethylbenzene radical cation but relatively slow from hexaethylbenzene. Aromatic carbonyl compounds undergo fairly efficient nuclear acetamidation at a Pt anode in moist MeCN/Et4NBF4 [63]. Alkenes undergo the combined addition of NHCOCH3 and F when oxidized in MeCN containing Et4NF,3HF [64] and of NHCOCH3 and PhSe when oxidized in MeCN containing PhSeSePh [65]. It is of interest to notice that while fluoroacetamidation of cyclic alkenes results predominantly in cis-addition, selenoacetamidation gives the trans-isomer as the major product. n-Alkanes can be oxidized in MeCN/R4NBF4 solutions (see Chapter 23). The observation of well-defined oxidation waves, the consumption of 2F, and the formation of acetamides at secondary carbon atoms supports the mechanism outlined in Equation 19.25 for n-octane [66]. In some cases, prior cleavage of carbon–carbon bonds is observed, again in a manner that suggests a carbocation intermediate, Equation 19.26 [67]:

(+ Isomers)

–2e––H+

CH3CN

+

H2O

(anodically generated halogen and thus position 3: 35%) (position 4: 32%) NHCOCH3

–H+ +N C CH3

33% CH3CN, H2O

NHCOCH3 +

–2e–, –2H+

NHCOCH3

(19.25)

(19.26)

In the wake of the aforementioned early work, several investigations of anodic acetamidation of polycyclic hydrocarbons were initiated; substituted adamantanes were much used in this work because of their availability and their relative ease of oxidation [68–70]. The results are generally explicable in terms of expected stabilities of intermediate carbocations and competition between proton loss and C–C bond cleavage. Difunctionalization of adamantanes is also possible using anodic acetamidation [70]: R

NHCOCH3

E = 2.5 V vs. Ag/Ag+ CH3CN, H2O

R + NHCOCH3

E = 3.0 V vs. Ag/Ag+ as above

(R = H: 74%) (R = tert-Bu: 62%)

NHCOCH3 (R = H: 58%) NHCOCH3

© 2016 by Taylor & Francis Group, LLC

(19.27)

783

Oxidative Substitution and Addition Reactions

Esters, ketones, and alcohols may also undergo substitution in MeCN [71–73]. Some of these substrates are oxidized in a region where oxidation of the electrolyte is concurrent and, certainly for esters, it may well be that an indirect mechanism operates with necessary initial abstraction of hydrogen by a radical derived from the electrolyte [71]. Examples are given in Equations 19.28 and 19.29. In one case, the anodic oxidation of 2-octanone in MeCN/LiClO4, the reaction mechanism has unambiguously been demonstrated to involve specific abstraction of one hydrogen atom at C-5 by the carbonyl oxygen in a six-membered ring cyclic transition state, reminiscent of the first chemical step in the gas-phase McLafferty rearrangement [73]. NHCOCH3 CH3CN, H2O

(19.28)

–2e–, –2H+

O

O

NHCOCH3 CH3CN, H2O

OCH3

OCH3

–2e–, –2H+

O

OCH3

+

O

(19.29)

NHCOCH3 O

Acetamidation is also a result of the anodic oxidation in MeCN solution of carboxylates (see Chapter 33) and alkyl halides [68,74–77]. Much discussion and experimentation has focused on the mechanism of cleavage of the alkyl halides, particularly the iodides. From the extent of rearrangement of cations supposedly produced by alkyl iodide oxidation compared with that for the same cations produced by solvolysis of tosylates, it has been suggested that in electrolysis, MeCN assists cation formation [74]. The SN2 character of this process receives support from the 20% inversion observed for oxidation of optically active 2-iodooctane, Equation 19.30. An alternative mechanism must, however, be considered that is based on the known anodic oxidation, at the potentials used for alkyl iodide cleavage, of iodine to an iodine (I) species [75]. The process given in Equation 19.31, which is catalytic in iodine, is therefore plausible, given the initial generation of some iodine by direct cleavage: Pt anode, CH3CN

(-) I

+ 3- and 4-Isomers NHCOCH3 70% (75% Racemization)

R I + I+

I2 + R+

CH3CN, H2O

30%

R NHCOCH3

–H+

(19.30)

(19.31)

Anode

A Ritter-type reaction is also observed during the anodic oxidation of phenylthiomethane derivatives [78] of the type shown in Equation 19.32, which summarizes the proposed reaction mechanism: OH

SPh

OH

SPh

R

+

R –e–

O

O

O

OH

–SPh

O

(19.32) OH

+ R O

© 2016 by Taylor & Francis Group, LLC

O

NHCOCH3

CH3CN, H2O –H+

R O

O

784

B.

Organic Electrochemistry

N-CYANOMETHYLATION AND AMINATION

The solvent, MeCN, may be incorporated in the product in an altogether different type of process, which leads to formation of substituted acetonitriles. For instance, the oxidation of 2,2,6,6-tetramethylpiperidine, or the corresponding morpholine, in MeCN/NaClO4 under oxygen-free conditions leads to the formation of an aminoacetonitrile derivative [79]. The reaction was suggested to follow the mechanism given in the following [X = CH2,O]:

X

NH

X

N

+

X

N

+

–e–

+ NH

X

CH3CN

CH2CN

X

–H+

N

X

NH + CH2CN

X

N CH2CN

(19.33)

Essentially, the same mechanism has been proposed for anodic α-amination of tetrahydrofuran, which may be accomplished by oxidation of, for example, R2NH or R2NLi, in this solvent [80]. The best yields (70–80%) were observed when R2NH was piperidine.

C.

NITRATION

The formation of nitro compounds upon electrolysis of a suspension of an aromatic hydrocarbon in dilute aqueous HNO3 has been known for a long time [81]. It has been interpreted in terms of the formation of a high concentration of HNO3 near the anode surface, in which region an ordinary electrophilic nitration process would take place. Studies on the anodic oxidation of NO3– in nitromethane [82] would seem to confirm this assumption; N2O5 is the oxidation product, probably formed via reactions (19.34) and (19.35). In aqueous solution, this would lead to the formation of HNO3. 2 NO3–

–4e–

NO2+ + NO3–

2 NO2+ + O2

(19.34)

N2O5

(19.35)

Similarly, the oxidation of N2O4 in, for example, concentrated HNO3 or H2SO4 [83], MeCN [84], or sulfolane [85] has been shown to lead to nitrating mixtures containing NO2+. Another mechanism of anodic nitration involves the generation of a radical cation in the presence of NO2, Equation 19.36. This mechanism has attracted considerable attention [86–88] in the vivid discussion of electron transfer as a general step in conventional nitration, where in Equation 19.36, the nitronium ion instead of the anode serves to oxidize the substrate to the radical cation state in the first step. Here, it is sufficient to notice that, using naphthalene and methyl-substituted naphthalene radical cations as substrates in CH2Cl2/Bu4NPF6, it was shown [86] that the coupling reaction between ArH• + and NO2 is more selective than nitration by acetyl nitrate in (CH3CO)2O or N2O4 in CH2Cl2.

© 2016 by Taylor & Francis Group, LLC

785

Oxidative Substitution and Addition Reactions + ArH

–e–

ArH

NO2

+ H Ar NO2

–H+

ArNO2

(19.36)

The anodic oxidation of solutions of aromatic hydrocarbons and AgNO2 in MeCN gives mononitration with a substitution pattern that is held to support initial formation of a radical cation [89]. Alkenes may similarly be oxidized in the presence of AgNO2, and depending on the alkene, nitro products are formed by substitution or addition. In probably the best example, coelectrolysis of anthracene and AgNO2 at 1.3 V versus SCE gave 9-nitroanthracene in 47% yield (current efficiency 28%), and from 1,1-diphenylethylene, electrolyzed at 1.5 V, 1,1-diphenyl-2-nitroethylene was obtained in 40% yield. In each case, 1F was consumed. However, since NO2– is more easily oxidized than both anthracene and 1,1-diphenylethylene, it cannot be excluded that NO2– is oxidized at the anode to give NO2/N2O4, in itself a good nitrating reagent in homogeneous medium [86]. Constant potential electrolysis of 1,4-dimethoxybenzene in micellar solutions containing NO2– showed, however, that the nitration product, 2,5-dimethoxynitrobenzene, was only formed at potentials sufficiently high to oxidize both NO2– and 1,4-dimethoxybenzene. This was taken as evidence in support of a reaction including the coupling between the radical cation and NO2 [90]. The same type of mechanism was proposed for the electrochemical nitration of naphthalene in the presence of NO2– in aqueous nonionic surfactant solutions [91]. Aliphatic gem-dinitro compounds can be prepared by anodic oxidation of a nitroparaffin in aqueous alkaline solution in the presence of an excess of NO2– [92]. Yields are considerably better when Ag+ is the oxidizing agent [93], so that an indirect oxidation process involving the generation of Ag+ at an Ag anode was found to give excellent results [94]. Electrochemical nitration has been applied also to functionalize single-wall carbon nanotubes [95] and for the introduction of the nitro group in horse heart myoglobin [96].

D. CYANATE ION AS NUCLEOPHILE Both nuclear and side-chain substitutions of alkylaromatics by NCO– have been observed [97]. Thus, electrolysis of anisole in MeOH/KOCN results in substitution of hydrogen by a methylcarbamoyl group, via the isocyanate. The reaction has been pictured as a direct oxidation process. Yields are fairly low ( CN– ≈ OH– > Cl– ≈ CH3O – > Br –. The mechanistic possibilities are given in the following: CH3OH

RH

+

–e–, –H+ X

CH2OH

or

CH3O–

R

HX

(X

+

R

R

–e–

–e–

CH3O

= CH2OH or CH3O )

(19.48) (19.49) (19.50)

R+

+ CH3O

and/or

(19.51)

ROCH3

R+ + CH3OH

–H+

One would not expect F– or ClO4 – to be oxidized under the conditions used, which suggests that MeOH might be the electroactive species, forming a radical intermediate that then functions as X• • in Equation 19.49. The radical CH2OH is the most likely candidate since the O–H bond in MeOH is much stronger than the C–H bond. In alkaline medium, MeO – is probably the electroactive species, and MeO• would then be X•. Allylic methoxylation is often the result of anodic oxidation of alkenes and dienes in methanolic media [127,128]. Cyclopropanes behave similarly [129]. The reactions are usually accompanied by products of addition and combination as shown in Scheme 19.1 although even more complex reaction mixtures may result. The same is true for the methoxylation of styrenes [130] and stilbenes [131]. The yields for allylic methoxylation are only moderate (ca. 25%) and compare unfavorably with those for the analogous allylic acetoxylation (see following text) [132]. The methoxylated products that result from substitution of aromatic hydrocarbons are as the rule more easily oxidized than the starting materials, and for that reason, the reaction usually leads to mixtures of di- and trimethoxylated products as observed, for example, for naphthalene [133]. Exceptions are the methoxylation of anthracene [134] that results in the substitution–addition product (~20%), Equation 19.52, and of acenaphthylene [135] that results in the addition product (60–80%, mixture of cis- and trans-isomers), Equation 19.53: H3CO

OCH3

CH3OH

(19.52)

–6e–, –6H+ H3CO

© 2016 by Taylor & Francis Group, LLC

OCH3

790

Organic Electrochemistry R2

R1 –e– R1 R2

R2

x2 –2H+

R1

R1

R1

–H+

–e–

2CH3OH

R1

OCH3 R2

R2

R2

CH3OH –H+

x2

CH3O

R2

+

x2

R1

OCH3

R2 +

R1

R2 CH3OH

R1 CH3OH –e–, –H+

–H+

R1

R2 OCH3

OCH3 R2

R1

OCH3 (+ Isomers)

SCHEME 19.1

Oxidation of alkenes in methanol—products and pathways. H3CO

OCH3

(19.53)

CH3OH –2e–, –2H+

As mentioned earlier, the methoxylation of alkylbenzenes to the side-chain-substituted products appears to proceed via in an indirect mechanism, and it is (still) not easy to predict the combination of substrate concentration, conversion, and supporting electrolyte that result in the maximum yield of product. This is illustrated by a careful study [136] of the methoxylation of 1,3-diisopropylbenzene that with potassium bromate as supporting electrolyte gives 1,3-bis(2-methoxypropan-2-yl)benzene in 80% isolated yield: CH3OH

H3CO

OCH3

(19.54)

KBrO3, i = 40 mA

Attempted direct anodic methoxylation of substituted 2-methyl and 2-benzylnaphthalenes results in dimerization for electron-rich derivatives and in nuclear substitution, in the 6-position, for electronpoor derivatives. Only traces of side-chain oxidation were found [137]. However, indirect oxidation, Equation 19.55, with DDQ mediation results in clean conversion (70%) and on a preparative scale into the corresponding carboxaldehyde or ketone [137]. The conversion of 2-methoxy-6methylnaphthalene into the carboxaldehyde is important because it is potentially an intermediate for the preparation of the nonsteroidal anti-inflammatory drug naproxen. CH3 CH3O

Undivided cell, graphite electodes constant current aq. HOAc, Et4NOTs, 80°C DDQ

–2e–, –2 H+ DDQH2

© 2016 by Taylor & Francis Group, LLC

CHO CH3O

70%

(19.55)

791

Oxidative Substitution and Addition Reactions

The products of methoxylation of phenols are usually further oxidized under the reaction conditions used [104,138]. Experimental conditions can be found, however, which in some cases allow selective formation of para-methoxylated dienones, ortho-methoxylated dienones, or dimeric products. It is believed that methoxylation involves attack of MeOH on anodically generated phenoxylium ions, whereas dimeric products arise by combination of phenoxyl radicals: OH

Dimers

O

–e–, –H+

O

O

O

H –e–

+

CH3OH

OCH3

+

–H+ H

(19.56)

OCH3

The methoxylation of anisoles [139–150] and other alkoxybenzenes [146], methoxynaphthalenes [139], and methoxyanthracenes [143,147] has received considerable attention. For the oxidation of alkylanisoles in MeOH containing various electrolytes [140], the current–potential curves support the scheme given, for p-methylanisole in the following: +OCH 3

CH3O

OCH3

H3C

OCH3

CH3OH + OCH3

OCH3

–H+ H3C CH3OH –e–, –H+

OCH3

(19.57)

–e– CH3

–H+, –e–

CH3

OCH3

OCH3 CH3OH –H+

CH+2

CH2OCH3

The oxidation of 1,4-dimethoxybenzene in MeOH to yield the corresponding quinone bisketal, Equation 19.58, is the prototype example of the numerous related 1,4-addition reactions that may be accomplished electrochemically [139–148]: OCH3

CH3O

OCH3

2 CH3OH

(19.58)

–2e–, –2H+ OCH3

CH3O

OCH3

Thioanisoles behave differently, and usually sulfonium ions result from oxidation. However, substitution reactions are favored in the presence of a silica-supported base and when electronwithdrawing groups are present in the α-position [151]: 0.1 M NaClO4, CH3OH Base

S CHF2

© 2016 by Taylor & Francis Group, LLC

Undivided cell

OCH3 S CHF2

(19.59)

792

Organic Electrochemistry

Additions similar to those reported for methoxy-substituted aromatic compounds are observed also for vinyl ethers, Equation 19.60 [152,153], inden-1-one [154], furans [155,156], thiophenes [157], and pyrroles [158]: RO

RO

2CH3OH

(19.60)

–2e–, –2H+ CH O 3

OCH3

In contrast, methoxylation of N,N-dimethylaniline in MeOH/KOH leads predominantly to sidechain substitution [159,160] most likely according to the following mechanism: H3C

+ N(CH3)2

N(CH3)2

N

H3C

CH2

N

CH2+

H3C

N

CH2OCH3

(19.61)

CH3OH –e–

–e–

–H+

–H+

N,N-Dimethylbenzylamine [161,162], 2,4,6-tris(dialkylamino)-1,3,5-triazines [163], N-(2,2,2trifluoroethyl)amines [164,165], and N-(2,2-difluoroethyl)amines [165] exhibit similar behavior. Methoxylation of enamines results in substitution at both the vinylic and the allylic positions [153], whereas methoxylation of imines and the related oxazolines gives rise to α-substituted products [166]. Oxidation of anilides such as N-benzyl-4-methylaniline in 5% aqueous methanol in the presence of sodium bicarbonate affords 4-methoxy-4-methylbenzoquinol N-benzoylimine and a dimer, 4-[N-benzoyl-N-(4-methylphenyl)amino]-4-methylbenzoquinol N-benzoylimine [167]: O H– N

O

O

NaHCO3 LiClO4 CH3OH

N

N +

O H3C

CH3

OCH3

H3C

(19.62) N

CH3

Another reaction that has attracted considerable attention is the alkoxylation of N,N-dialkylamides and related compounds [168–184]. A typical example is the oxidation of N,N-dimethylamides in ROH/NH4NO3 [169]: O CH3 R1 C N CH3

R2OH –2e–, –2H+

O CH3 R1 C N CH2OR2

R1 = H, CH3, C6H5 R2 = CH3, CH3CH2, CH3CH2CH2CH2

(19.63)

The use of a nitrate salt results in better yields and cleaner reaction than is observed with NaOR. A mechanism involving initial discharge of NO3– to form NO3•, which then abstracts a hydrogen atom from one of the N-alkyl groups, is considered the most likely one in this case. When, for example, NaOMe is used as supporting electrolyte, the reaction most likely involves the substrate radical cation as the primary intermediate followed by deprotonation and oxidation to the corresponding carbocation, which finally reacts with methanol. In this context, it is of interest to notice that the intramolecular alkoxylation expected for the series of hydroxyamides shown in Equation 19.64, n = 0−4, did not take place in MeCN; only polymeric material was formed [170]. When the reaction

© 2016 by Taylor & Francis Group, LLC

793

Oxidative Substitution and Addition Reactions

was instead carried out in MeOH, only the normal methoxylation products were formed, Equation 19.65. These, however, could easily be converted to the cyclized products by addition of acid. H3C

O N C

H3C n

HO H3C

O N C

O N C

O H3C –2e–, –2H+

(19.64)

H2C

–2e–, –2H+

CH3OH

H3C HO

H3C

CH3CN

n

O N C

(19.65)

CH3OCH2

n

n

HO

Numerous examples of the application of this synthetically useful reaction have been reported. Illustrative examples include the synthesis of eneamides and enecarbamates [171,172] and a large series of pyrrolidine, piperidine, and perhydroazepine derivatives [171,173–176]. Attempts to methoxylate smaller ring systems like N-acetylaziridine and N-formylazetidine failed and resulted in ring-opened products [177]. Methoxylation lends itself well to operation on a relatively large scale [178–182]. For instance, several N-formyl derivatives have been anodically methoxylated on a 0.5 kg scale using a capillary gap cell with a graphite anode [179,181]. An example is given in Equation 19.66. Polyalkoxylation of N,N-dimethylformamide has also been achieved [180]. O

O CH3OH

N CHO

(92%)

8.4 F

(19.66)

N OCH3 CHO

The selectivity and stereochemistry of the anodic methoxylation of the cyclic amides are noteworthy. First, N-formyl-2-methylpiperidine gives the 6-methoxy-substituted product in spite of the fact that loss of the proton in the 2-position would give a tertiary carbocation, which intuitively would be expected to be the more stable intermediate. Second, 4-substituted N-formylpiperidines give the corresponding 2-methoxy derivative with the methoxy group occupying an axial position [184]. Both these experimental facts were attributed to the steric constraints imposed by the N-formyl group. A good example of the selectivity is shown in Equation 19.67; the methoxylated product could be cyclized to the bicyclic amide by reaction with TiCl4 in CH2Cl2 [183]: O

R N

O

R n

CH3OH, CH3CN

N

–2e–, –2H+

O

R n

TiCl4, CH2Cl2

N

n

(19.67)

Cl

OCH3

Cyclic amides carrying a carboxylic acid group in the α-position to the nitrogen atom have been observed to undergo a non-Kolbe-type substitution reaction by oxidation in methanol [185,186], for example, with silica-supported piperidine as the base [185]: Si H3C

N H3C

O N

CH3OH COOH 4F

© 2016 by Taylor & Francis Group, LLC

(19.68)

O N

OCH3

794

Organic Electrochemistry

The anodic methoxylation of amides is not restricted to simple formic and AcOH derivatives, but has been performed also with peptides, Equation 19.69, [187], N-substituted carbamates [188–190], and cyclic derivatives such as N-(methoxycarbonyl)pyrrolidines [135c,190–192], N-methoxycarbonylpiperidines [190,193,194], and oxazolidin-2-ones [195–197]:

Boc

H N

H N

O CH3OH

OCH3

N H

Boc

–2e–, –2H+

O

OCH3 N H

O

OCH3

(19.69)

O

The prototype reaction [188] for methoxylation of a carbamate is shown in Equation 19.70, and the applicability is further illustrated by the examples given in Equations 19.71 [191], 19.72 [194a], and 19.73 [196]. It is seen that the stereochemical configuration of the starting material is preserved in all cases. O CH3 RO C N CH3

–2e–, –2H+

CH3OH

COOCH3

N

CH3OH

–2e–, –2H+

O CH3 RO C N CH2OCH3

(19.70)

CH3O

(19.71)

COOCH3 CH3COO

COOCH3

N

COOCH3 CH3COO

OCOCH3

–2e–, –2H+

N

OCOCH3

CH3OH CH3O

COOCH3

COOCH3 O HN

(19.72)

N

O O

CH3OH

HN

O

trans/cis = 8/1

–2e–, –2H+ CH3O

CH2OH

(19.73)

CH2OH

The presence of α-silyl [198] or α-phenylthio [199] groups activates carbamates toward oxidation by 500–600 mV and results in regioselective introduction of a methoxy group at the α-carbon carrying the activating substituent:

N

O C

OCH3

N

CH3OH 2–3 F

X

O C

OCH3

X = Si(CH3)3, SPh

(19.74)

OCH3

In all the examples given so far, the substrate carries at least one N-α-hydrogen atom. The anodic oxidation of fully substituted amides, like N,N-di-tert-butylformamide and N-formyl-2,2,6,6tetramethylpiperidine, in MeOH would be expected to follow a different pathway. The products isolated after 12–14F, methyl N-tert-butylcarbamate and N-methoxycarbonyl-2,2,6,6-tetramethylpiperidine, respectively, Equation 19.75 [200], indicated that the primarily formed substrate radical cation loses the formyl proton. Further oxidation of the neutral radical leads to the cation, which may either undergo cleavage, Equation 19.76, or nucleophilic attack by the solvent, Equation 19.77: (CH3)3C N CHO (CH3)3C

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–2e–, –H+

(CH3)3C + N C O (CH3)3C

(19.75)

795

Oxidative Substitution and Addition Reactions (CH3)3C N C O (CH3)3C + N C O (CH3)3C

2CH3OH +

C+

CH3OH (CH3)3C (CH3)3C

(CH3)3C OCH3

–H

(CH3)3

(CH3)3C NHCOOCH3

(19.76) (19.77)

NCOOCH3

Sulfonamides [201,202], sulfinylamines [203], and amidophosphates [201,204] have been used as starting materials as well. Saturated ethers, for example, 1,4-dioxane, can be α-methoxylated by electrolysis in MeOH/ NaOMe or MeOH/NH4NO3 in 10–30% yield, Equation 19.78 [205]. Yields are considerably higher, 45–70%, for the heterocyclic 1,3-benzodioxoles [206]. O O

O

CH3OH –2e–, –2H+

(19.78) O

OCH3

Related to these reactions is the anodic methoxylation of sulfides carrying an electron-withdrawing substituent, for example, cyano [207], fluoromethyl [208], difluoromethyl [208], or trifluoromethyl [208,209] in the α-position, Equation 19.79, and the fluoride-mediated methoxylation of thiazolidines, 1,3-oxathiolanes and 1,3-dithiolanes [210]. An analogous reaction has been reported for selenides carrying electron-withdrawing perfluoroalkyl or cyano substituents [211]. Ar

S CH2 X

CH3OH –2e–, –2H+

Ar

S CH X OCH3

X = CN, CH2F, CHF2, CF3

(19.79)

In analogy with what has been observed for carbamates, the trimethylsilyl substituent in, for example, 1-phenylthio-1-trimethylsilylalkanes is easily replaced by methoxy during anodic oxidation in MeOH [212,213]. Similarly, the anodic oxidation of α,α′-bis(trimethylsilyl)xylenes in MeOH results in replacement of one trimethylsilyl group by methoxy [214].

C. ACYLOXYLATION (ACETOXYLATION) Substitution by an acyloxy group can be accomplished by electrolysis of a variety of organic compounds in the presence of a carboxylate in a suitable solvent. The most commonly encountered process is acetoxylation that takes place in AcOH or AcOH/Ac2O. The supporting electrolyte in this case normally is NaOAc, although Me4NNO3, Bu4NBF4, NaClO4, and others can be employed in special cases. Other acyloxylations using carboxylates whose corresponding acids cannot be used as solvent may be performed in DMF or MeCN. For trifluoroacetoxylation (see following text), a variety of relatively deactivated aromatic substrates or aliphatic compounds may be employed. Acetoxylation of alkenes takes place in the allylic position [132,215,216] concurrent with addition of acetoxy groups across the double bond or double-bond system. The mechanism is almost certainly analogous to that shown in Scheme 19.1 for methoxylation. Because AcO – is difficult to oxidize, the anodic discharge of simple, unactivated alkenes may be achieved in its presence. Subsequent reactions of the cationic intermediates are not very selective, but several experimental parameters may be controlled, and the method can compare well with the few competing chemical methods. The oxidation of n-alkanes is possible in CF3COOH containing R4NBF4 with the carbocations so formed being trapped to give secondary trifluoroacetates [66,217,218]. Anodic trifluoroacetoxylation of long-chain aliphatic ketones [219] or carboxylic acids [217] gives rise to mixtures of

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796

Organic Electrochemistry

positional isomers. The substitution pattern for the ketones indicates that the reaction proceeds via intramolecular abstraction of a hydrogen atom from the 5-position by the carbonyl group. This reaction has been observed to take place during the oxidation of aliphatic ketones in MeCN (see previous text). The trifluoroacetoxylation of carboxylic acids is less selective, and the product distribution gives little indication of the mechanism. The yields of acetoxylation products for aromatic and alkylaromatic compounds vary greatly [3,7–10,220–222], depending on the oxidation potential of the product(s) formed. As a general rule, they tend to be low in nuclear acetoxylation of aromatic compounds and fair to good in acetoxylations at activated C atoms. A characteristic feature of the reaction is that nuclear acetoxylation of aromatic compounds requires the presence of AcO – [223], whereas acetoxylation of the other types of substrates takes place with either AcO – or other anions present [132,169,224]. Both nuclear and side-chain acetoxylation may be rationalized according to the general mechanism given in the following: CH2+

CH2

CH2OCOCH3 CH3

–e–

+

CH3

CH3

COO–

–H+

–e–

CH3

CH3

CH3

(19.80)

CH3COO– +

–e– H

OCOCH3

H

–H+ OCOCH3

OCOCH3

(+ ortho isomer)

In the context of the direct mechanism, several perplexing features concerning nuclear and side-chain substitution of aromatic compounds have been highlighted [225], and in the search for explanations, attempts have been made to estimate, by thermochemical calculations, free energies of activation for reactions between radical cations and nucleophiles. Long-standing puzzles in this area include the dependence of the ratio of side-chain-/nuclear-substituted products on the nucleophile; for alkylbenzenes, side-chain acetoxylation predominates in AcOH unless AcO – is present, when nuclear substitution becomes important. The isomer distribution for anodic acetoxylation of a number of monosubstituted benzenes has been determined [223]. The reaction closely resembles ordinary electrophilic aromatic substitution processes, perhaps on the side of low-selectivity reactions. The deuterium kinetic isotope effect, k H/k D, for nuclear acetoxylation in anisole was found to be 1.0, whereas for α-substitution in ethylbenzene, a value of 2.6 was observed. The interpretation of these values is not straightforward [225]. From a preparative point of view, it is of interest to notice that selective nuclear substitution may be obtained when the oxidation is carried out in an undivided cell using a cathode made of Pd/C [220]. With this experimental setup, the side-chain acetoxylated product is reduced back to starting material at the cathode. The nuclear-substituted product is unaffected and will accumulate during the electrolysis. Acetoxylation of 1,4-dimethoxybenzene results in the ortho-substituted product [226]: OCH3

OCH3 OCOCH3

COO–

CH3

–2e–, –2H+ OCH3

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(19.81) OCH3

797

Oxidative Substitution and Addition Reactions

Aromatic compounds, in particular those carrying electron-withdrawing substituents, may be oxidized in the presence of CF3COO − with consequent trifluoroacetoxylation [63,227–234]. For compounds PhX, X may be (CH3)3C, COOMe, NO2, CF3, COMe, COPh, and CN; the trifluoroacetoxy derivatives produced by anodic oxidation in CH2Cl2/CF3COOH/Et4NBF4 [224], CF3COOH/ CF3COONa [228], or CH3NO2/CF3COOH [229] are hydrolyzed during work-up to phenols. Yields of the phenols so produced are in the region 20–80%, and usually, substantial amounts of the meta-substituted phenols are formed. For chlorobenzene, prolonged electrolysis in CF3COOH/ CF3COONa, followed by hydrolysis, leads to good yields of the corresponding resorcinol and pyrocatechol derivatives [230]. A number of intramolecular acyloxylations have been reported [235–243], two examples of which are given in the following: CH2CH2COO–

HO

–O

O

–2e–, –H+

(19.82)

O O

O

CH3OH –2e–, –H+

O

(19.83)

O

CH3O

The α-acyloxylation of N-alkyl-substituted amides has received attention [169,174,244] from both mechanistic and preparative points of view. In electrolytes containing alkali metal carboxylates, the direct mechanism probably operates, whereas in case where a nitrate salt is the supporting electrolyte, an indirect mechanism involving hydrogen abstraction from the N-alkyl group by anodically generated NO3• is indicated. Similarly, 1-(p-tolylsulfonyl)azetidine is converted to the corresponding 2-acetoxy derivative by anodic oxidation in AcOH/AcONa, [245]: CH3COOH CH3COONa

N Ts

–2e–, –H+

OCOCH3

(19.84)

N Ts

Acyloxylation of a number of benzyl ethers has been reported, Equation 19.85 [246]. A hemiacetal acetate is presumably formed and subsequently hydrolyzed during the work-up procedure to give the aldehyde.

Ar

CH2OCH3

–2e–, –H+

Ar

CH3COO– + CHOCH3 Ar

OCOCH3 CH OCH3

H+, H2O

Ar

CHO

(19.85)

30–70%

Aryl alkyl sulfides and selenides containing an electron-withdrawing substituent such as trifluoromethyl in the α- or β-position may be α-acetoxylated in good-to-excellent yields [209,247–252]. In the absence of the electron-withdrawing substituent, the yields are usually poor [209,247] and give a mixture of the S-oxide and the α-acetoxylated product [252]. However, if the oxidation is carried out in boiling AcOH/Ac2O/AcONa, the corresponding α-acetoxy derivatives may be obtained in almost quantitative yields [252]. Under these conditions, the S-oxide, if formed at all, is converted to the desired α-acetoxy derivative in a Pummerer reaction parallel to the direct electrochemical acetoxylation. This illustrates that reactions that at room temperature give rise to mixtures of products may be synthetically useful when carried out at higher temperatures at which the thermodynamically more stable products are formed.

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798

D.

Organic Electrochemistry

NITRATE ION AS NUCLEOPHILE

Concurrent with side-chain acetoxylation of alkylaromatic compounds in AcOH/NH4NO3, the formation of benzyl nitrates, ArCH2ONO2, is observed in about the same yield as the ArCH2OAc [224,253,254]. It seems likely that NO3− and AcOH compete as nucleophiles for an intermediate benzyl cation, although an indirect mechanism cannot be ruled out completely because of the complexity of the supporting electrolyte behavior [255].

V. FORMATION OF CARbON-TO-SULFUR bONDS A.

THIOLS AS NUCLEOPHILES

Thiols are easier to oxidize than most organic substrates and can for that reason usually not be used as nucleophiles in anodic substitution processes. An exception is the cooxidation of catechols and 2-mercaptopyrimidines that leads to fair-to-good yields of the corresponding substitution products in reactions that most likely include the ortho-benzoquinones as intermediates [256]: R3 R1 R2

OH

–2e–, –2H+

OH

B.

N

R1 R2

O

R1

R3

SH N

N N

O

R2

OH

S

OH

(19.86)

THIOCYANATION

Phenols [6c], anisoles, Equation 19.87 [257], and aromatic amines [6c] are thiocyanated on electrolysis in aqueous or acetic acid solutions containing SCN–. The reaction is indirect since there is compelling evidence that the thiocyanating reagent must be (SCN)2, formed by anodic oxidation of SCN– [258]. In MeCN solution, oxidation of SCN– and SeCN– in the presence of phenol or various aromatic amines leads to thiocyanation and selenocyanation in 55–80% yields [259]. OCH3

CH3COOH, SCN–

OCH3 (~80%)

1.5–2.2 F

(19.87)

SCN

The direct reaction between a radical cation derived from an aromatic hydrocarbon (naphthalene) and SCN− has been studied by mixing the salt (C10H8)2PF6 with a solution of Bu4NSCN in CH2Cl2 at −78°C [260]. Two major products were isolated: 1-naphthylthiocyanate (kinetic control, 28%) and 2-naphthylisothiocyanate (thermodynamic control, 8%). The mechanistic details leading to this product distribution are not clear.

VI. FORMATION OF CARbON-TO-HALOgEN bONDS A.

FLUORINATION

The virtually unique suitability of electrochemical methodology for preparing fluorinated compounds warrants the inclusion of a separate treatment on this subject (Chapter 20).

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799

Oxidative Substitution and Addition Reactions

B.

CHLORINATION AND BROMINATION

The oxidation of chloride or bromide ions to the atomic or molecular halogen is relatively easy; consequently, the mechanism of anodic halogenation, brought about by oxidation of organic substrates in the presence of chloride [261–269] or bromide [270–274], is not always clear. In many cases, it may involve halogenation by anodically generated halogen and thus, resembles conventional organic halogenations. In other cases, where the substrate is easily oxidized, the halide plays the role of a nucleophile and attacks a radical cation. Examples include the chlorination of alkenes [262,269] and of aromatic hydrocarbons [261,268] and substituted derivatives such as 1,4-dimethoxy-2-tertbutylbenzene [265] and aniline [263] as well as substituted azulenes, Equation 19.88 [267]. Anodic chlorination has been used also for the postfunctionalization of poly(3-hexylthiophene) [275]. COOCH3

COOCH3 + Cl–

(19.88)

–2e–, –H+ Cl

Similarly, successful brominations have been reported for, for instance, alkenes [273,274], aromatic hydrocarbons [270], acetanilide [271], and acetylated d-glycals [272]: OAc

OAc +2 Br–

O

AcO

–2e–

AcO

Br

O

AcO AcO

OAc +

O

AcO AcO Br

Br

(19.89)

Br

9:1

C.

IODINATION

The oxidation of I2 in MeCN produces a very reactive although ill-defined iodine(I) species [75,276]. The production of this species, in either MeCN or CH2Cl2 in the presence of a variety of substituted benzenes, results [276] in efficient iodination, for example, as shown in the following: OCH3 ½ I2

CH3CN –e–

+ “H3C C N I”

I

–H+

OCH3

(19.90)

90%

Formation of a C–I bond is also the result of anodic oxidation of, for example, iodobenzene in the presence of benzene [277–279]. Two reactions are in competition, Equation 19.91; a careful investigation of the electrode kinetics supports the mechanism given and also allows optimization of conditions to the point where a 95% yield of the diphenyliodonium cation is obtained [278]. Hydrolysis of this species completes a route from benzene to phenol. PhI

+

or PhI

Dimeric products

+ I

–e–

I

(19.91)

PhI –e–, –H+ + I

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H2O

OH

800

Organic Electrochemistry

Electrochemical oxidation of terminal alkynes in MeOH in the presence of NaI as supporting electrolyte has been reported to result in the formation of the corresponding 1-iodo derivatives in 70–80% yield [280]: R

H

CH3OH, NaI –2e–

R

I

(19.92)

REFERENCES 1. (a) Hammerich, O.; Parker, V.D. Adv. Phys. Org. Chem. 1984, 20, 55; (b) Workentin, M.S.; Parker, V.D.; Morkin, T.L.; Wayner, D.D.M. J. Phys. Chem. A 1998, 102, 6503 and references cited therein; (c) Schmittel, M.; Ghorai, M.K. In Electron Transfer in Chemistry, Matthay, J. (ed.), vol. 2, chap. 1, WileyVCH: Weinheim, Germany, 2001. 2. (a) Eberson, L. Electron Transfer Reactions in Organic Chemistry, Springer-Verlag: Berlin, Germany, 1987 and references cited therein; (b) Adv. Phys. Org. Chem. 1982, 18, 79; (c) J. Am. Chem. Soc. 1983, 105, 3192; (d) Eberson, L.; Jönsson, L.; Sänneskog, O. Acta Chem. Scand. 1985, B39, 113; (e) Walling, C.; Zhao, C., El-Taliawi, G.M. J. Org. Chem. 1983, 48, 4910. 3. (a) Baciocchi, E.; Eberson, L.; Rol, C. J. Org. Chem. 1982, 47, 5106; (b) Baciocchi, E.; Bartoli, D.; Rol, C.; Ruzziconi, R.; Sebastiani, G.V. J. Org. Chem. 1986, 51, 3587. 4. (a) Jönsson, L.; Wistrand, L.G. J. Chem. Soc., Perkin Trans. I 1979, 669; (b) Nyberg, K.; Wistrand, L.G. J. Org. Chem. 1978, 43, 2613. 5. Sheldon, R.A.; Kochi, J.K. Metal-Catalyzed Oxidations of Organic Compounds, Academic Press: New York, 1981. 6. (a) Techniques of Chemistry, Vol. 5, Parts 1 and 2: Techniques of Electroorganic Synthesis, N.L. Weinberg (ed.), Wiley-Interscience: New York, 1974; (b) Techniques of Chemistry, vol. 5, Part 3: Techniques of Electroorganic Synthesis, N.L. Weinberg and B.V. Tilak, (eds.), Wiley-Interscience: New York, 1982; (c) Weinberg, N.L.; Weinberg, H.R. Chem. Rev. 1968, 68, 449. 7. (a) Shono, T. Tetrahedron 1984, 40, 811; (b) Electroorganic Chemistry as a New Tool in Organic Synthesis, Springer: Berlin, Germany, 1984; (c) Electroorganic Synthesis, Academic Press: London, U.K., 1991. 8. Yoshida, K. Electrooxidation in Organic Chemistry, Wiley-Interscience: New York, 1984. 9. Torii, S. Electroorganic Syntheses, Part I: Oxidations, VCH: Weinheim, Germany, 1985. 10. (a) Eberson, L.; Nyberg, K. Adv. Phys. Org. Chem. 1976, 12, 1; (b) Tetrahedron 1976, 32, 2185; (c) Petrosyan, V.A. Mendeleev Commun. 2011, 21, 115. 11. (a) Nyberg, K. Acta Chem. Scand. 1970, B24, 1609; (b) Eberson, L.; Nyberg, K.; Sternerup, H. Acta Chem. Scand. 1973, B27, 1679. 12. Eberson, L.; Nyberg, K.; Sternerup, H. Chem. Scripta 1973, 3, 12. 13. Ronlán, A.; Bechgaard, K.; Parker, V.D. Acta Chem. Scand. 1973, B27, 2375. 14. Svanholm, U.; Parker, V.D. J. Am. Chem. Soc. 1976, 98, 2942. 15. Tommasino, J.-B.; Brondex, A.; Médebielle, M.; Thomalla, M.; Langlois, B.R.; Billard, T. Synlett 2002, 1697. 16. (a) Koyama, K.; Susuki, T.; Tsutsumi, S. Tetrahedron Lett. 1965, 627; (b) Koyama, K.; Susuki, T.; Tsutsumi, S. Tetrahedron 1967, 23, 2675; (c) Susuki, T.; Koyama, K.; Omori, A.; Tsutsumi, S. Bull. Chem. Soc. Jpn. 1968, 41, 2663. 17. Parker, V.D.; Burgert, B.E. Tetrahedron Lett. 1965, 4065. 18. Tsutsumi, S.; Koyama, K. Discuss. Faraday Soc. 1968, 45, 247. 19. Yoshida, K.; Fueno, T.; J. Chem. Soc., Chem. Commun. 1970, 711. 20. Eberson, L.; Nilsson, S. Discuss. Faraday Soc. 1968, 45, 242. 21. Andreades, S.; Zahnow, E.W. J. Am. Chem. Soc. 1969, 91, 4181. 22. Parker, V.D.; Eberson, L. J. Chem. Soc., Chem. Commun. 1972, 441. 23. (a) Eberson, L.; Helgée, B. Chem. Scripta 1974, 5, 47; (b) Acta Chem. Scand. 1975, B29, 451; (c) Acta Chem. Scand. 1977, B31, 813; (d) Acta Chem. Scand. 1978, B32, 313. 24. Yoshida, K. J. Am. Chem. Soc. 1977, 99, 6111. 25. Yoshida, K.; Fueno, T. Bull. Chem. Soc. Jpn. 1978, 60, 229 and references cited therein. 26. Atobe, M.; Aoyagi, T.; Fuchigami, T.; Nonaka, T. Electrochemistry 2004, 72, 821. 27. Liu, W.; Ma, Y.; Yin, Y.; Zhao, Y. J. Heterocycl. Chem. 2006, 43, 681. 28. Yoshida, K.; Toyo-oka, Y.; Takeda, K. J. Heterocycl. Chem. 1995, 32, 701. 29. Yoshida, K.; Kitabayashi, H. Bull. Chem. Soc. Jpn. 1987, 60, 3693. 30. Yoshida, K.; Takeda, K.; Minagawa, K. J. Chem. Soc., Perkin Trans. I 1991, 1119.

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Oxidative Substitution and Addition Reactions 31. 32. 33. 34. 35. 36. 37. 38. 39. 40.

41. 42.

43. 44. 45. 46. 47. 48. 49.

50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. 61. 62. 63. 64. 65. 66. 67. 68. 69. 70.

801

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193. Shono, T., Matsumura, Y., Fujita, T. Tetrahedron Lett. 1991, 32, 6723. 194. (a) Plehiers, M., Hootelé, C. Can. J. Chem. 1996, 74, 2444; (b) Durant, A.; Hootelé, C. Can. J. Chem. 1992, 70, 272; (c) Driessens, F.; Hootelé, C. Can. J. Chem. 1991, 69, 211. 195. Martre, A.M.; Mousset, G.; Prudhomme, M.; Rodrigues-Pereira, E. Electrochim. Acta 1995, 40, 1805. 196. Danielmeier, K.; Schierle, K.; Steckhan, E. Tetrahedron 1996, 52, 9743. 197. Wang, P.-C. Heterocycles 1985, 23, 2237. 198. Yoshida, J.; Isoe, S. Tetrahedron Lett. 1987, 28, 6621. 199. Sugawara, M.; Mori, K.; Yoshida, J. Electrochim. Acta 1997, 42, 1995. 200. Blum, Z.; Nyberg, K. Acta Chem. Scand. 1981, B35, 743. 201. Shono, T.; Matsumura, Y.; Tsubata, K.; Uchida, K.; Kanazawa, T.; Tsuda, K. J. Org. Chem. 1984, 49, 3711. 202. Bodmann, K.; Bug, T.; Steinbeisser, S.; Kreuder, R.; Reiser, O. Tetrahedron Lett. 2006, 47, 2061. 203. Turcaud, S.; Martens, T.; Sierecki, E.; Pérard-Viret, J.; Royer, J. Tetrahedron Lett. 2005, 46, 5131. 204. Sierecki, E.; Turcaud, S.; Martens, T.; Royer, J. Synthesis 2006, 3199–3208. 205. Shono, T., Matsumura, Y. J. Am. Chem. Soc. 1969, 91, 2803. 206. Thomas, H.G.; Schmitz, A. Synthesis 1985, 31. 207. Kimura, M.; Koie, K.; Matsubara, S.; Sawaki, Y.; Iwamura, H. J. Chem. Soc., Chem. Comm. 1987, 122. 208. Fuchigami, T.; Yano, H.; Konno, A. J. Org. Chem. 1991, 56, 6731. 209. (a) Fuchigami, T.; Nakagawa, Y.; Nonaka, T. Tetrahedron Lett. 1986, 27, 3869; (b) Fuchigami, T., Yamamoto, K.; Konno, A. Tetrahedron 1991, 47, 625; (c) Fuchigami, T.; Yamamoto, K.; Nakagawa, Y. J.  Org. Chem. 1991, 56, 137; (d) Tajima, T.; Fuchigami, T. J. Am. Chem. Soc. 2005, 127, 2848; (e) Tajima, T.; Fuchigami, T. Chem. Eur. J. 2005, 11, 6192. 210. Baba, D.; Fuchigami, T. Electrochim. Acta 2003, 48, 755. 211. Surowiec, K., Fuchigami, T. Tetrahedron Lett. 1992, 33, 1065. 212. Jouikov, V.; Fattahova, D. Electrochim. Acta 1998, 43, 1811. 213. Yoshida, J.; Matsunaga, S.; Murata, T., Isoe, S. Tetrahedron 1991, 47, 615. 214. Koizumi, T.; Fuchigami, T., Nonaka, T. Electrochim. Acta 1988, 33, 1635. 215. Adams, C.; Frankel, E.N.; Utley, J.H.P. J. Chem. Soc., Perkin Trans. I 1979, 353. 216. Kowalski, J.; Płoszyńska, J.; Sobkowiak, A.; Morzycki, J.W.; Wilczewska, A.Z. J. Electroanal. Chem. 2005, 585, 275–280. 217. (a) Hembrock, A.; Schäfer, H.J.; Zimmermann, G. Angew. Chem. Int. Ed. Engl. 1985, 24, 1055; (b) Schäfer, H.J.; Cramer, E.; Hembrock, A.; Matusczyk, G. Electroorganic Synthesis, R.D. Little, and N.L. Weinberg (eds.), Marcel Dekker: New York, 1991, p. 169. 218. Koch, V.R.; Miller, L.L. J. Am. Chem. Soc. 1973, 95, 8631. 219. Campbell, C.B.; Pletcher, D. Electrochim. Acta 1978, 23, 923. 220. Eberson, L.; Oberrauch, E. Acta Chem. Scand. 1981, B35, 193. 221. Eberson, L.; Webber, A. Acta Chem. Scand. 1982, B36, 53. 222. Pei, J.; Qin, S.; Li, G.; Hu, C.; Chin. J. Chem. Phys. 2011, 24, 244. 223. Eberson, L. J. Am. Chem. Soc. 1967, 89, 4669. 224. Ross, S.D.; Finkelstein, M.; Petersen, R.C. J. Am. Chem. Soc. 1967, 89, 4088. 225. (a) Eberson, L.; Blum, Z.; Helgée, B.; Nyberg, K. Tetrahedron 1978, 34, 731; (b) Eberson, L.; Nyberg, K. Acta Chem. Scand. 1978, B32, 235; (c) Eberson, L., Jönsson, L., Wistrand, L.-G. Acta Chem. Scand. 1978, B32, 520. 226. (a) Petrosyan, V.A.; Vakhotina, T.S.; Burasov, A.V. Russ. Chem. Bull. Int. Ed. 2005, 54, 1580; (b) Burasov, A.V.; Petrosyan, V.A. Mendeleev Commun. 2008, 18, 196. 227. Svanholm, U.; Parker, V.D. Tetrahedron Lett. 1972, 471. 228. Blum, Z.; Cedheim, L.; Nyberg, K., Eberson, L. Acta Chem. Scand. 1975, B29, 715. 229. So, Y.-H.; Miller, L.L. Synthesis 1976, 468. 230. Bockmair G.; Fritz, H.P.; Gebauer, H. Electrochim. Acta 1978, 23, 21. 231. Fritz, H.P.; Kremer, H.J. Z. Naturforsch. 1976, 31b, 1565. 232. Kreh, R.P.; Tadros, M.E.; Hand, H.M.; Cockerham, M.P.; Smith, E.K. J. Appl. Electrochem. 1986, 16, 440. 233. (a) Fujimoto, K.; Maekawa, H.; Tokuda, Y.; Matsubara, Y.; Mizuno, T.; Nishiguchi, I. Synlett 1995, 661; (b) Fujimoto, K.; Tokuda, Y.; Maekawa, H., Matsubara, Y., Mizuno, T.; Nishiguchi, I. Tetrahedron 1996, 52, 3889. 234. Utley, J.H.P.; Elinson, M.; Güllü, M.; Ludwig, R.; Motevalli, M. Acta Chem. Scand. 1999, 53, 901. 235. (a) Eberson, L.; Nyberg, K. J. Am. Chem. Soc. 1966, 88, 1686; (b) Acta Chem. Scand. 1964, 18, 1568. 236. Iwasaki, H.; Cohen, L.A.; Witkop, B. J. Am. Chem. Soc. 1963, 85, 3701. 237. Scott, A.I.; Dodson, P.A.; McCapra, F.; Meyers, M.B. J. Am. Chem. Soc. 1963, 85, 3702.

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Organic Electrochemistry Bonner, W.A.; Mango, F.D. J. Org. Chem. 1964, 29, 430. Koehl, W.J. J. Org. Chem. 1967, 32, 614. Banda, F.M.; Brettle, R. J. Chem. Soc., Perkin Trans. I 1974, 1907. Laurent, A.; Laurent, E.; Thomalla, M. C. R. Acad. Sci. Ser. C 1972, 274, 1537. Myall, C.J.; Pletcher, D.; Smith, C.Z. J. Chem. Soc., Perkin Trans. I 1976, 2035. Adams, C.; Jacobsen, N.; Utley, J.H.P. J. Chem. Soc., Perkin Trans. II 1978, 1071. Rand, L.; Mohar, F. J. Org. Chem. 1965, 30, 3156; 3885. Shono, T.; Matsumura, Y.; Uchida, K.; Nakatani, F. Bull. Chem. Soc. Jpn. 1988, 61, 3029. Garwood, R.F.; Naser-ud-din; Weedon, B.C.L. J. Chem. Soc., Chem. Commun. 1968, 923. Surowiec, K.; Fuchigami, T. J. Org. Chem. 1992, 57, 5781. Fuchigami, T.; Yamamoto, K.; Yano, H. J. Org. Chem. 1992, 57, 2946. Jouikov, V.; Ivkov, V.; Fattahova, D. Tetrahedron Lett. 1993, 34, 6045. Tajima, T.; Fuchigami, T. Angew. Chem. Ind. Ed. 2005, 44, 4760. Tajima, T.; Kishi, Y.; Nakajima, A. Electrochim. Acta 2009, 54, 5959. Almdal, K.; Hammerich, O. Sulfur Lett. 1984, 2, 1. Ross, S.D.; Finkelstein, M.; Petersen, R.C. J. Org. Chem. 1970, 35, 781. Nyberg, K. Acta Chem. Scand. 1970, B24, 473. Rao, R.R.; Milliken, S.B.; Robinson, S.L.; Mann, C.K. Anal. Chem. 1970, 42, 1076. Zeng, C.; Ping, D.; Zhang, S.; Zhong, R.; Becker, J.Y. J. Electroanal. Chem. 2008, 622, 90. Gitkis, A.; Becker, J.Y. Electrochim. Acta 2010, 55, 5854–5859. Cauquis, G.; Pierre, G. C. R. Acad. Sci., Ser. C 1968, 266, 883. Cauquis, G.; Pierre, G. C. R. Acad. Sci., Ser. C 1971, 272, 609. Fritz, H.P.; Ecker, P. Chem. Ber. 1981, 114, 3643. Forsyth, S.R.; Pletcher, D.; Healy, K.P. J. Appl. Electrochem. 1987, 17, 905 and references cited therein. Faita, G.; Fleischmann, M.; Pletcher, D. J. Electroanal. Chem. 1970, 25, 455. Matsuda, Y.; Nishiki, T.; Sakota, N.; Nakagawa, K. Electrochim. Acta 1984, 29, 35. Verniette, M.; Pouillen, P.; Martinet, P. Bull. Soc. Chim. Fr. 1981, I-343. Appelbaum, L.; Danovich, D.; Lazanes, G.; Michman, M.; Oron, M. J. Electrochem. Chem. 2001, 499, 39. Stevanović, D.; Damljanović, I.; Vukićević, M.; Manojlović, N.; Radulović, N.S.; Vukićević, R.D. Helv. Chim. Acta 2011, 96, 1406. Ungureanu, E.M.; Razus, A.C.; Birzan, L.; Buica, G.; Cretu, M.; Cristian, E. Electrochim. Acta 2006, 52, 794. Raju, T.; Kulangiappar, K.; Kulandainathan, M.A.; Shankar, G.K.; Muthukumaran, A. Electrochim. Acta 2005, 51, 356. Milisavljević, S.; Vukićević, R.D. J. Serb. Chem. Soc. 2004, 69, 941. Millington, J.P. J. Chem. Soc. B 1969, 982. Abirami, D.; Chithra, B.; Krishanamoorthy, T.K. Asian J. Chem. 2010, 22, 834. (a) Damljanović, I.; Vukićević, M.; Manojlović, D.; Sojic, N.; Buriez, O.; Vukićević, R.D.; Electrochim. Acta 2010, 55, 965; (b) Čolović, M.; Vukićević, M.; Šegan, D.; Manojlović, D.; Sojic, N.; Somsák, L.; Vukićević, R.D. Adv. Synth. Catal. 2008, 350, 29. Milisavljević, S.S.; Wurst, K.; Laus, G.; Vukićević, M.D.; Vukićević, R.D. Steroids, 2005, 70, 867. Ogamino, T.; Mori, K.; Yamamura, S.; Nishiyama, S. Electrochim. Acta 2004, 49, 4865. Hayashi, S.; Inagi, S.; Hosaka, K.; Fuchigami, T. Synth. Metals 2009, 159, 1792. Miller, L.L.; Watkins, B.F. J. Am. Chem. Soc. 1976, 98, 1515. Miller, L.L.; Hoffmann, A.K. J. Am. Chem. Soc. 1967, 89, 593. Hoffelner, H.; Lorch, H.W.; Wendt, H. J. Electroanal. Chem. 1975, 66, 183. Peacock, M.J.; Pletcher, D. J. Electrochem. Soc. 2001, 148, D37. Nishiguchi, I.; Kanbe, O.; Maekawa, H. Synlett 2000, 89–91.

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20

Fluorination Toshio Fuchigami and Shinsuke Inagi

CONTENTS I. II.

Introduction ...........................................................................................................................807 Selective Electrochemical Fluorination in Organic Solvents ................................................808 A. Anodic Fluorination of Aromatic Compounds..............................................................808 B. Anodic Benzylic Fluorination .......................................................................................808 C. Anodic Fluorination of Olefins ......................................................................................809 D. Anodic Fluorination of Aldehydes ................................................................................ 810 E. Anodic Fluorination of Organosulfur Compounds ....................................................... 810 F. Anodic Fluorination of Other Heteroatom Compounds ................................................ 811 G. Anodic Fluorination of Heterocyclic Compounds......................................................... 812 1. Anodic Fluorination of Heterocyclic Rings Having a Phenylthio Electroauxiliary .... 812 2. Anodic Fluorination at the Side Chain of Heterocyclic Sulfides ............................ 813 3. Anodic Fluorination of Heterocyclic Ring ............................................................. 814 III. Anodic Fluorination in Ionic Liquids.................................................................................... 816 A. Solvent-Free Anodic Fluorination ................................................................................. 816 B. Anodic Fluorination in Ionic Liquids Under Ultrasonication ....................................... 819 C. Double Ionic Liquid System for Anodic Fluorination ................................................... 820 D. Effects of Ethereal Additives on Anodic Fluorination .................................................. 820 IV. Anodic Fluorination of Macromolecules .............................................................................. 821 V. Indirect Anodic Fluorination ................................................................................................ 821 VI. Anodic Fluorination Using Alkali Metal Salts ..................................................................... 823 VII. Conclusion ............................................................................................................................. 823 References ......................................................................................................................................824

I. INTRODUCTION Organofluorine compounds are classified into two groups: perfluoro-compounds and partially fluorinated compounds. The compounds in the former class are widely utilized as functional materials, while those in the latter family find biological uses such as pharmaceuticals and agrochemicals. Perfluoro-compounds are manufactured by converting all C–H bonds to C–F bonds using electrochemical fluorination in anhydrous liquid HF as a solvent with a nickel anode. This process was developed more than 60 years ago. Later, electrochemical perfluorination in a KF–2HF melt at a carbon anode was developed for the preparation of perfluorinated low-molecular-weight organic compounds [1]. Both of these processes are now used commercially. In contrast, selective electrochemical fluorination is a rather new field and methodology, but it has not been well developed because of low reaction selectivity, the low nucleophilicity of fluoride ions, and competitive anode passivation (the formation of a nonconducting polymer film on the anode surface that suppresses faradaic current). The fluorination can be commonly achieved in aprotic solvents (acetonitrile, dichloromethane, dimethoxyethane [DME], nitromethane, sulfolane, etc.) containing fluoride ions to provide mostly mono- and/or difluorinated products [1–4]. Electrolyses are conducted at constant potentials slightly higher than the first oxidation potential of

807

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808

Organic Electrochemistry

the substrate by using a platinum or graphite anode. Constant current electrolysis is also effective for selective fluorination in many cases. The choice of the combination of a supporting fluoride salt and an electrolytic solvent is most important to accomplish efficient selective fluorination because competitive anode passivation takes place very often during the electrolysis. Pulse electrolysis is in many cases effective in order to avoid such passivation. Therefore, difficult-to-oxidize fluoride salts, which do not cause the passivation of the anode and have strongly nucleophilic F−, are generally recommended as the supporting fluoride salts. Thus, room temperature molten salts such as R3NnHF (n = 3–5), R4NF-nHF (n = 3–5), and 70% HF/pyridine (Olah's reagent) are most often used and even R4NBF4 and R4NPF6 salts are effective in some cases [1–4]. Particularly when HF supporting salts and low hydrogen overpotential cathodes such as platinum are used, the reduction of protons (hydrogen evolution) occurs predominantly at the cathode during the electrolysis. Therefore, a divided cell is not always necessary for the fluorination under such conditions. In aprotic solvent, F− becomes more nucleophilic; however, the reactivity of F− is quite sensitive to the water content of the electrolysis system because a hydrated F− is a weak nucleophile. Drying of both the solvent and electrolyte is therefore necessary to optimize the formation of fluorinated products. Since the discharge potential of the fluoride ion is extremely high (>+2.9 V vs. SCE [saturated calomel electrode] at Pt anode in acetonitrile [MeCN]), the fluorination proceeds via a (radical) cation intermediate as shown in Equation 20.1, which is the general pathway for anodic nucleophilic substitutions. F R

F– RH

–e

RH+

F

–e

+R

–H+

H

H

–H+

F– R

R F

(20.1)

–e +R

II.

SELECTIVE ELECTROCHEMICAL FLUORINATION IN ORgANIC SOLVENTS

A.

ANODIC FLUORINATION OF AROMATIC COMPOUNDS

In 1970, Rozhkov et al. reported the first example of selective electrochemical fluorination of aromatic compounds such as benzene and naphthalene in MeCN containing the ionic liquid Et4NF3HF (Equation 20.2) [5]. Since that time, selective electrochemical fluorination of various aromatic compounds such as benzene, substituted benzenes, anthracene, and 9,10-diphenylanthracene has been accomplished similarly by constant potential anodic oxidation [1,2,6–10]. F –2e, –H

F

+

(20.2)

+

Et4NF-3HF/MeCN 1.8 V vs. SCE 27%

3%

F

B. ANODIC BENzYLIC FLUORINATION Generally, anodic benzylic substitution reactions take place readily. However, anodic benzylic fluorination does not always occur. The major competitive reaction is acetamidation when MeCN is used as a solvent. Laurent et al. found that anodic benzylic mono- and difluorination proceeds selectively when the benzylic position is substituted by electron-withdrawing groups (EWGs) (Equation 20.3) [11]. In these cases, p-methoxy or p-chloro substituents on the benzene are necessary for the operation of efficient fluorination. In their absence, benzylic acetamidation becomes a major reaction (Equation 20.4).

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809

Fluorination –2e, –H+ p-XC6H4CH2-EWG

Et3N-3HF/MeCN

–2e,–H+ p-XC6H4CHF-EWG Et3N-3HF/MeCN (36–73%)

p-XC6H4CF2-EWG (48–95%)

(20.3)

X = MeO, Cl EWG = COAr, COOEt, CN, SO3Et

Ar

–2e, –H+

COMe

Ar

COMe

Ar

COMe

+

F–/MeCN

F

Ar = Ph: Ar = p-MeOC6H4:

NHCOMe

7% 69%

(20.4)

34% ethylbenzene > toluene [12]. On the other hand, the efficiency of acetamidation is reverse. Moreover, triphenylmethane is selectively monofluorinated to provide fluorotriphenylmethane in high yield (80%) even in MeCN [13]. These facts suggest that the more stable benzylic cation intermediate reacts with fluoride ion more efficiently.

C. ANODIC FLUORINATION OF OLEFINS Anodic oxidation of olefins in the presence of fluoride ions provides mono- and/or difluorinated products (Equation 20.5) [14–16]. R R1

–2e Et3N-3HF/MeCN or AcOH

R

Y

F

R1

(20.5)

Y = F, NHCOMe, OAc

Anodic fluorination of vinyl sulfides such as 2-(phenylthio)styrene provides vicinal difluorides [17]. 1-Phenylhexene undergoes stereoselective difluorination and fluoroacetamidation upon anodic oxidation in MeCN, while the difluorination predominates in the less nucleophilic solvent, dichloromethane (CH2Cl2) [18]. Recently, it was shown that nitromethane (MeNO2) significantly promoted electrochemical fluoro-selenenylation and iodofluorination of electron-deficient olefins such as α,β-unsaturated esters, amides, and phosphonates (Equation 20.6) [19,20]. Me

Me

–2e Et3N-5HF/MeNO2

PhSeSePh or n-Bu4NI + COOEt

F

Y COOEt

(20.6)

Y = PhSe (quant.), l (75%)

Anodic fluorination of α-acetoxystyrene and 1-acetoxy-3,4-dihydronaphthalene provides the corresponding α-fluoroketones (Equation 20.7) [21,22]. OCOMe R

F –2e Et3N-3HF

R = H, Me

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OCOMe R F

O R aq. NaHCO3

F

(20.7)

810

Organic Electrochemistry

Yoshida et al. achieved anodic fluorination accompanied by carbon–carbon bond formation using silyl, stannyl, or thio groups as electroauxiliaries (EAs) as shown in Equation 20.8 [23–25]. EA –2e Bu4NBF4/CH2Cl2

O C7H15

O C7H15

F

(20.8)

50–98%

EA = MeS, Bu3Sn, Me3Si

D. ANODIC FLUORINATION OF ALDEHYDES Selective anodic formyl hydrogen-exchange fluorination of aliphatic aldehydes in Et3N-5HF/MeCN provides the corresponding acyl fluorides in good yields as shown in Equation 20.9 [26]. In these reactions, fluoride salts like Et3N-3HF lead to lower yields, reflecting the fact that this salt is discharged prior to oxidation of the carbonyl group. In the case of aromatic aldehydes, the aromatic ring is also fluorinated simultaneously. O R C H

–2e Et3N-5HF/MeCN 2.4–2.6 V vs. Ag/Ag+

O R C F

(20.9)

R = 3-Heptyl: 89% R = Cyclohexyl: 84% R = tert-Butyl: 66%

E. ANODIC FLUORINATION OF ORGANOSULFUR COMPOUNDS Fuchigami et al. and Laurent et al. independently found that anodic fluorination of sulfides is markedly promoted by α-electron-withdrawing groups (EWGs) such as fluoroalkyl, ester, and cyano groups to provide the corresponding α-fluorinated products in good yields as shown in Equation 20.10 [27–31]. This was the first successful example of selective electrochemical fluorination of chalcogen compounds. The anodic fluorination is highly regioselective and widely applicable, even though α-thio-substituted esters have several positions susceptible to substitution by a fluorine. A fluorine is introduced exclusively α to the ester group, and no fluorination of the p-tolyl, benzyl, or heptyl groups is observed (Equation 20.10) [31]. As shown in Equation 20.10, the electron-withdrawing ability of the fluoroalkyl group does not affect the efficiency of anodic monofluorination of sulfides. Even simple alkyl phenyl sulfides devoid of an EWG undergo fluorination in ethereal solvents such as tetrahydrofuran (THF) and DME to provide monofluorinated products in moderate yields [30]. R1

S

R2

–2e, –H+ Et3N-3HF/MeCN

F R1

S

R2

(20.10)

~88% R1 = Ph, R2 = CF3 (62%); CF2H (53%); CFH2 (69%) R1 = Ph, p-MeC6H4, PhCH2, C7H15 R2 = CH3, C CH, COOEt, CN, COMe, COPh, CONH2, PO(OEt)2

The fluorination proceeds by way of a Pummerer-type mechanism via the fluorosulfonium cation (A), as shown in Equation 20.11 [28,30]. Thus, when R is an EWG, the deprotonation of A is significantly facilitated, and consequently, the fluorination proceeds efficiently.

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811

Fluorination F– H + PhS CHR F

–2e F–

PhS CH2R

F–

+

PhS CHFR

PhS CHR

–HF

(20.11)

A

Since then, several groups have studied the selective electrochemical fluorination of various heteroatom compounds including heterocyclic compounds in organic solvents as follows. Simonet et al. similarly performed regioselective anodic α-monofluorination in an Et3N-3HF/MeCN solution of alkyl aryl sulfides, having an EWG on the aromatic ring [32]. Interestingly, anodic fluorination of α-thio α-arylesters provides fluorodesulfurization products while that of α-thioacids provides fluorinated products accompanying decarboxylation as shown in Equation 20.12 [33]. Me p-XC6H4

C

Me

–2e, –PhS+ COOEt

F– (R = Et)

F 55–84%

p-XC6H4

C

Me

–2e,–H+, –CO2

COOR

p-XC6H4

F– (R = H)

SPh

C SPh

F

(20.12)

65–70% X = Cl, MeO, Me2CHCH2

F. ANODIC FLUORINATION OF OTHER HETEROATOM COMPOUNDS Anodic α-monofluorination of selenides bearing α-electron-withdrawing cyano and ester groups can be performed in Et3N-3HF/MeCN using an undivided cell, while the fluorination of α-selenoacetoamide requires an anion-exchange membrane diaphragm (Equation 20.13) [34]. –2e, –H+ PhSeCH2R

Et3N-3HF/MeCN

PhSeCHFR

(20.13)

R = CN, COOEt, CONH2 (60–70%)

In contrast, anodic oxidation of organotellurium compounds in the presence of fluoride ions results in difluorination at the tellurium atom predominantly in excellent yields and with high current efficiencies (Equation 20.14) [35]. In this case, α-fluorination does not occur. F Ph Te

R

–2e Ph

Et3N-3HF

Te

R

F 75–86%

R = Me, CHF2, CH2CF3, PH

(20.14)

Anodic oxidation of tetra-alkylsilanes in the presence of fluoride ions provides the corresponding fluorosilanes derived from cleavage of the C–Si bond [36]. Becker and Shakkour found that anodic oxidation of cyclic peralkylsilanes results in the formation of α,ω-difluorosilanes via Si–Si bond cleavage (Equation 20.15) [37]. Pr Pr Pr Si –2e Si Si Pr Pr Et 4NBF4 Si Si Pr Pr Pr Pr Pr

Pr

Pr

Pr Pr Pr + Si Si Si Si Pr Pr Pr Pr F F F F Pr

Si

30%

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54%

(20.15)

812

Organic Electrochemistry

Anodic fluorination of ethers such as DME and diethyleneglycol dimethylether results in monofluorination at the terminal carbon selectively, while that of crown ethers undergoes carbon–carbon bond cleavage preferentially on fluorination to provide the α,ω-difluorinated products in good yields (Equation 20.16) [38,39]. O

–2e, –H+

O

O

O

F–

n

O

O

n n = 0 (62%), 1 (55%)

F

(20.16)

Anodic fluorodeiodination of alkyl iodides provides the corresponding alkyl fluorides chemoselectively (Equation 20.17) [40]. –e, –1/2 l2

R–l

F–

(20.17)

R–F 72–85%

R = Me(CH2)3, AcO(CH2)10, Cl(CH2)10, MeCO(CH2)10

Anodic oxidation of benzophenone hydrazone in Et3N-3HF/CH2Cl2 gives mainly diphenylmonofluoromethane (Equation 20.18) [41]. Ph

F

–ne

NNH2

Et3N-3HF/CH2Cl2

Ph

H

Ph

F +

Ph 95%

F

Ph

Ph

(20.18)

3%

Anodic oxidation of organic compounds containing group 15 elements in the presence of fluoride ions provides the corresponding fluorinated products (Equation 20.19) [42]. Fluorination occurs at the heteroatoms selectively. Ph

Ph Y Ph

Ph

–2e Et3N-3HF

F Y

Ph

F Ph

(20.19)

Y = P, Sb

G. ANODIC FLUORINATION OF HETEROCYCLIC COMPOUNDS Partially fluorinated heterocycles very often show unique and pronounced biological activities, they are essential for development of new types of agrochemicals and medicines. However, limited examples of selective anodic fluorination of heterocycles have been reported [43–46]. These processes are limited to only nitrogen- and oxygen-containing heterocycles, and the yields are generally quite low. In addition, no successful anodic fluorination of sulfur-containing heterocycles has been reported until 1991. Fuchigami et al. have developed conditions for highly selective anodic fluorination reactions of various heterocyclic compounds [3,4]. 1. Anodic Fluorination of Heterocyclic Rings Having a Phenylthio Electroauxiliary Heterocyclic compounds having a phenylthio group as an electroauxiliary are selectively oxidized to result in regioselective α-fluorination. Thus, various α-phenylthio lactones and lactams including β-lactams can be anodically fluorinated efficiently Equations 20.20 and 20.21 [47,48] (SSCE = saturated sodium calomel electrode). Cyclic phosphonate can be also similarly fluorinated as shown in Equation 20.22 [49].

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813

Fluorination O

O SPh

X

–2e, –H+ Et3N-3HF

SPh F

X

n

(20.20)

n

n = 1, X = O (84%) n = 0,1,2, X = NR (69–92%)

F

SPh N R

Et3N-3HF/MeCN 1.8–2.0 V vs. SSCE 2.3–4.0 F

O

ArS

SPh

–2e, –H+

O

N R

R = Et, i-Pr, n-Bu, t-Bu, c-Hexyl, Bn Yield: 65–92%

–2e, –H+

P O O

(20.21)

O

F

SAr O P O O

Et4NF-3HF/DME

(20.22)

58%

Highly regioselective, anodic monofluorination of oxindole and 3-oxo-1,2,3,4-tetrahydroisoquinoline can be achieved by using Et4NF-mHF (m = 3, 4) (Equation 20.23) [50]. Although three kinds of benzylic carbons exist in a molecule (n = 1 in Equation 20.23), the fluorination takes place at the 4-position exclusively. F

SPh O N

n R

–2e, –H+ Et4NF-mHF/MeCN

SPh O

(20.23)

N

n R

n = 0, R = Ph, p-Tol: 50–64% (m = 4) n = 1, R = PhCH2: 71% (m = 3)

2. Anodic Fluorination at the Side Chain of Heterocyclic Sulfides Anodic monofluorination at the side chain of various heterocyclic compounds has been systematically studied [51,52]. The active methylenethio group attached to heterocycles is selectively fluorinated to give the corresponding α-fluorinated products (Equation 20.24). Notably, the use of MeCN prevented the formation of fluorinated products, while the use of DME markedly increased their yields [51–56]. The pronounced solvent effect of DME could be explained in terms of the significantly enhanced nucleophilicity of fluoride ions, as well as the suppression of anode passivation and overoxidation of fluorinated products. O

O N

Me

N

COPh S

–2e, –H+ COPh

Et4NF-4HF

N Me

N

COPh F S

COPh

in MeCN: 0% DME: 78%

© 2016 by Taylor & Francis Group, LLC

(20.24)

814

Organic Electrochemistry

Heterocyclic propargyl sulfides are also anodically fluorinated to give the α-fluoro products (Equation 20.25) [57]. In sharp contrast, in the case of thiazolyl sulfides and oxazolyl sulfides, the fluorination takes place on the heteroaromatic rings and no α-fluorinated products are formed (Equation 20.26) [58,59]. F Het

–2e, –H+

S

Het

F–

S

(20.25)

35–60%

Het = 2-Pyridyl, 4-Pyridyl, 2-Pyrimidyl, 2-Quinolyl, 2-Benzothiazolyl Me

Me F

–2e, –H+

N S

O

R

Et4NF-4HF

F

N S

O

R

(20.26)

55–70%

R = H, COMe, CN

3. Anodic Fluorination of Heterocyclic Ring Highly regioselective anodic monofluorination of 2-aryl-4-thiazolidinones can be performed by using pulse electrolysis in Et3N-3HF/MeCN (Equation 20.27) [60]. However, this electrolytic system is not suitable for anodic monofluorination of 2-substituted 1,3-dithiolan-4-ones and 1,3-oxathiolan-4-ones owing to severe passivation of the anode. In contrast, Et4NF-4HF provides monofluorinated products selectively (Equation 20.27) [61,62]. In these cases, benzylic fluorination does not take place at all although anodic benzylic substitution easily takes place in general. The high regioselectivity can be explained in terms of facilitation of deprotonation of radical cation intermediate of the substrates by the electron-withdrawing carbonyl group (i.e., kinetic acidity control). S

R X

F

–2e, –H+

R X

Et3NF-nHF/MeCN Yeild: 58–86%

O

S

(20.27)

O (trans major)

R = Ph, 2-Naphthyl, Mesityl, Et, n-Pr X = NH, MeN, i-PrN, PhN, BnN (n = 3) X = O,S (n = 4)

Diastereoselective electrochemical fluorination of various heterocyclic compounds derived from optically active amino acids and 1,2-diol has also been demonstrated, as depicted in Equations 20.28 and 20.29 [63–65]. Electrochemical fluorination of N-substituted pyrroles and its application to the synthesis of gem-difluorinated fused heterocyclic compounds has also been reported (Equation 20.30) [66]. Interestingly, dehydrodimers readily derived from benzothiazines underwent electrochemical fluorination accompanied by C–C double bond cleavage to provide gem-difluorinated benzothiazine derivatives, as shown in Equation 20.31 [67]. Quite recently, synthesis of fluorinated indole derivatives was also achieved by electrochemical fluorination of indoles [68]. COOMe S

NCOR

F –2e, –H+ Et3N-4HF/DME

R = p-Tol, Ph, Me, H

© 2016 by Taylor & Francis Group, LLC

S

COOMe NCOR

52–91% yield, 59–95% de

(20.28)

815

Fluorination

–2e, –H+ O

O

O

Et3N-3HF

O

SO2Ph F

SPh

SPh

B

OH

HO

HCI 50% aq. MeOH

m-CPBA

(20.29)

81% from B

F

in MeCN: 66% (80% de) DME/MeCN (1:1): 92% (61% de) Me N

CN

F

Water

–4e

Me N

F

Et3N–3HF/MeCN

F O

F N Me

(20.30)

O quant

54% Ar

N S

N

–4e

S

Et3N-3HF/DME Ar

Ar

2 S

N

F

(20.31)

F

A r = Ph: 67% p-BrC6H4: 62%

The selectivity of fluorination was strongly influenced by supporting fluoride salts (Equation 20.32) [69]. Since Et3N-3HF contains the free base Et3N, the difluorinated product, once formed, was dehydrofluorinated to the monofluoro product. O

O

O

H –2e

F O

Et4NF-4HF/MeCN

F Ph

F

–2e O

Et3N-3HF/MeCN

Ph

68% (cis/trans = 2)

O 58%

Ph

(20.32)

On the other hand, in the case of chroman-4-one derivatives, the fluorination does not take place at the olefin moiety but resulted at the α-position to the ring-oxygen atom to give the corresponding 2-fluorochromanones [70]. The same fluorinated product could also be obtained stereoselectively from an alternative anodic fluorination of homoisoflavone derivative, as shown in Equation 20.33 [70]. In these cases, Et4NF-4HF/DME is a suitable electrolytic solution. O

O

H –2e, –H+

Ar O

H

Et4NF-4HF/DME

Ar = Ph, p-ClC6H4, p-BrC6H4

O Ar

O

F 60–72%

CH2Ar

–2e, –H+ Et4NF-4HF/DME

O Ar = p-ClC6H4

(20.33) It is notable that DME also shows unique fluorination product selectivity (Equation 20.34) [71], which could be explained in terms of the stability of the anodically generated radical cation intermediate C (Equation 20.35). DME strongly stabilized the intermediate C, while CH2Cl2 destabilized C. Therefore, α-fluorination proceeded without desulfurization in DME, while fluorodesulfurization took place in CH2Cl2.

© 2016 by Taylor & Francis Group, LLC

816

Organic Electrochemistry F –e, –PhS

O

O

F

SAr F–/CH2Cl2

O

–2e, –H+

O

O

SAr

O

(20.34)

F–/DME

O

O

O

Ar = p-ClC6H4

SAr O

SAr –e

O

O

O

+

+

–e, –H+

O

O

O

SAr

O O

C Ar = p-ClC6H4 in CH2Cl2

–PhS

in DME F–

F–

F

F O

O

O

O

(20.35)

SAr

O O

III.

ANODIC FLUORINATION IN IONIC LIQUIDS

A.

SOLVENT-FREE ANODIC FLUORINATION

Solvent-free electrochemical fluorination is an alternative method for preventing anode passivation and acetoamidation. Handling extremely corrosive and poisonous anhydrous HF in a laboratory setting is accompanied by serious hazards and experimental difficulties. Ionic liquids such as 70% HF/ pyridine and commercially available Et3N-3HF [72] are often used to replace anhydrous HF. Meurs and Eilenberg first carried out solvent-free selective electrochemical fluorination, using the ionic liquid Et3N-3HF as the reaction medium, supporting electrolyte and fluorine source for anodic fluorination of benzenes, naphthalene, olefins, furan, benzofuran, and phenanthroline [73]. They obtained the corresponding partially fluorinated products in less than 50% yield (Equation 20.36). F F

F

F

–2e, –H+ N

N

(20.36)

Et3N-3HF

N

N 34%

Middleton et  al. similarly used the 70% HF/pyridine ionic liquid for anodic fluorination of 4-nitrotoluene and various compounds containing benzylic hydrogen atoms [74]. As shown in Equation 20.37, electrochemical fluorination of 4-fluorophenylacetonitrile in MeCN gave only 17% yield of α-fluorinated product, whereas a yield of 87% was obtained in the ionic liquid without MeCN. F

CH2CN

–2e, –H+ 70% HF/Py

F in MeCN: 17% 70% HF/Py: 87%

© 2016 by Taylor & Francis Group, LLC

CHFCN

(20.37)

817

Fluorination

Noel and Suryanarayanan have studied the voltammetry and electrochemical fluorination of PhSCH2CONH2 and PhCH2CN in Et3N-3HF [75]. They found that solvent-free Et3N-3HF had a much wider anodic potential window and obtained the desired α-fluorinated products in moderate to reasonable yields. However, Et3N-3HF and Olah's reagent are not anodically stable enough for certain purposes, as they are easily oxidized at around 2 V vs. Ag/Ag+. Momota et al. have developed a new series of ionic liquid fluoride salts with the general formula R4NF-nHF (n > 3.5, R = Me, Et, and n-Pr), that were useful in selective electrochemical fluorination [76]. These electrolytes were nonviscous liquids that had high conductivity and anodic stability. As a result, anodic partial fluorination of arenes such as benzene [76]; mono-, di-, and trifluorobenzenes [77]; chlorobenzene [78]; bromobenzene [79]; toluene [80,81]; and quinolones [82] was successfully carried out at high current densities using these ionic liquid fluoride salts in the absence of organic solvent with good to high current efficiencies (66–90%) (Equations 20.38 and 20.39). F

F F –2e–

(20.38)

Et4NF-4HF 2F

F F

F

90% CH2F

CH3 –2e, –H+

–2e, –H+

Et4NF-4HF 2.0 F 1.9 V vs. Ag/Ag+

Et4NF-4HF 2.0 F 2.1 V

52%

(20.39)

CHF2

CHF2 –2e, –H+ Et4NF-4HF 2.0 F 2.5 V

47%

F

Yoneda et al. have also investigated the electrochemical stability (potential window) of the ionic liquid Et3N-nHF (n = 3–5) by cyclic voltammetry [26]. They found that the anodic stability increased with increasing HF content (n) in the fluoride salts, while the cathodic stability showed the reverse tendency. Thus, Et3N-5HF was stable up to +3 V vs. Ag/Ag+, but was readily reduced at about −0.2 V to generate hydrogen gas. The potential window of Et4NF-4HF was almost the same as that of Et3N-5HF. Thus, the selective anodic fluorination of cyclic ketones and cyclic unsaturated esters in Et3N-5HF was successfully carried out to provide ring-opening and ring-expansion fluorinated products, respectively, as shown in Equations 20.40 and 20.41 [26,83]. O

R1 R2 n

O –2e

F

R1 R2

F

Et3N-5HF 2.0–2.4 V vs. Ag/Ag+

n

O NaOMe

MeO

R1 R2

F n

n = 0, R1 = R2 = H n = 0, R1 = Me, R2 = H n = 0, R1 = R2 = Me n = 1, R1 = R2 = Me

© 2016 by Taylor & Francis Group, LLC

26% 29% 91% 81%

(20.40)

818

Organic Electrochemistry F F CO2Et

–2e

CO2Et

Et3NF-5HF 2.2–2.3 V vs. Ag/Ag+ –20°C

n

(20.41)

n n=1 n=0

71% 56%

Noel et al. have employed Et3N-4HF as the electrolytic medium for electrochemical fluorination of N-alkylphenylacetamides [84] and indanone derivatives [85], as shown in Equation 20.42. In both cases, conversion and fluorinated product yield were moderate, and the products were a complex mixture. The potentiostatic conditions improved the monofluorination selectivity. O

O Et3N-4HF

O +

6F

F

F Constant current Constant potential

43.3% 69.8%

16.2% 7.3%

(20.42) O

O F F

+

+

F

+

Others

F 16.1% 0%

3.4% 0%

The fluorination of cyclic ethers, esters, lactones, and cyclic and acyclic carbonates can be achieved by anodic oxidation of a large amount of the liquid substrates and a small amount of Et4NF-4HF (only 1.5–1.7 equiv. of F− to the ether) at a high current density (150 mA/cm2) (Equations 20.43 through 20.45) [86,87]. X n O

–2e, –H+

X

Et3NF-4HF 2.0 F

O

n F

(20.43)

X = CH2, n = 0 83% X = O, n = 1 80% X = O, n = 0 77% O Z

O O

–2e, –H+

Z

O

Et4NF-5HF

(20.44)

F Z = CH2 Z=O O X

75% 87% O

–2e, –H+ OCH2CH3

Et4NF-5HF 2–2.5 F

X

OCHCH3 F

X = EtO X = Et

© 2016 by Taylor & Francis Group, LLC

97% 44%

(20.45)

819

Fluorination

Yoneda et al. have successfully carried out electrochemical fluorination of phenols in Et3N-5HF to provide 4,4-difluorocyclohexa-2,5-diene-1-ones, which were readily converted to p-fluorophenols in good yields by subsequent reduction with Zn (Equation 20.46) [88,89]. Hara and coworkers have investigated electrochemical fluorination of various phenols in Et3N-5HF and have found that carbon fiber cloth was a suitable anode, with various phenol derivatives being converted to 4,4-difluorocycohexadienone derivatives in good yields (Equations 20.47 and 20.48) [90]. O

OH t-Bu

t-Bu

OH t-Bu

t-Bu

+

–4e, –2H

t-Bu

Zn

t-Bu

(20.46)

H+

Et3N-5HF Pt anode

F

F

F 80%

80% (current efficiency) O

OH R

R

–4e, –2H+ Et3N-5HF Carbon fiber cloth

F

(20.47)

F

R = H: 61% t-Bu: 70% Ph: 77% O

OH +

–4e, –2H

Et3N-5HF R Carbon fiber cloth

(20.48)

R F F R = H: 89% Ph: 68%

The same group has also successfully carried out electrochemical fluorination of adamantanes in Et3N-5HF [91]. Mono-, di-, tri-, and tetrafluoroadamantanes were selectively prepared from adamantanes by controlling oxidation potentials, and the fluorine atoms were introduced selectively at the tertiary carbons, as shown in Equation 20.49. Adamantanes bearing functional groups such as ester, cyano, and acetoxymethyl moieties were also selectively fluorinated. –2e, –H+

–2e, –H+ F

Et3N-5HF 2.3 V vs. Ag/Ag+

F

Et3N-5HF 2.5 V

F

74%

79%

–2e, –H+ Et3N-5HF 2.7 V

(20.49)

F

F –2e, –H+ F F 61%

Et3N-5HF 3.0 V

F

F F 41%

B. ANODIC FLUORINATION IN IONIC LIQUIDS UNDER ULTRASONICATION The viscosity of ionic liquid fluoride salts is little higher than that of ordinary molecular solvents, which makes slow mass transport of substrate. However, Fuchigami et al. found that anodic fluorination of ethyl α-(phenylthio)acetate in Et3N-3HF proceeded smoothly without anode passivation

© 2016 by Taylor & Francis Group, LLC

820

Organic Electrochemistry

under ultrasonication to provide the monofluorinated product in high yield and with high current efficiency. Notably, anodic difluorination of the substrate could also be achieved in the same ionic liquid under ultrasonication, as shown in Equation 20.50 [92]. O SPh

EtO

O

–4e, –2H+

–2e, –H+ Et3N-3HF/MeCN 1.6 V vs. SSCE 2.5 F 76% yield

C.

F F

(20.50)

–2e, –H+ Et3N-3HF/MeCN 2.2 V vs. SSCE 20.7 F 53% yield

O SPh

EtO

SPh

EtO

6 F Et3N-3HF Under Ultrasonication 65% yield

F

DOUBLE IONIC LIQUID SYSTEM FOR ANODIC FLUORINATION

Fuchigami et al. have also found that a combination of Et4NF-nHF (n = 4, 5) and imidazolium ionic liquids was highly effective for the anodic fluorination of phthalides [93]. Since the oxidation potential of phthalide is extremely high (2.81 V vs. SCE), the yield is low even in ionic liquid fluoride salts due to simultaneous oxidation of the fluoride salts during electrolysis. In sharp contrast, when the ionic liquid 1-ethyl-3-methylimidazolium trifluoromethanesulfonate ([EMIM][OTf]) was used, the yield increased markedly as shown in Equation 20.51. The double ionic liquid system consisting of [EMIM][OTf] and Et3N-5HF enhanced not only the nucleophilicity of F− but also the electrophilicity of the cationic phthalide intermediate shown in Equation 20.51. The cationic intermediate generated from the phthalide was expected to have a TfO − counter anion (activated cation D in Equation 20.51), which readily reacted with F− to provide the fluorinated phthalide in good yield. – OTf –2e, –H+

+ O

O O

Et3N-nHF [C2mim][OTf ] n = 3 or 5

F F–

(20.51) O

O D Activated cation

D.

O

n=3 n=5

78% 90%

EFFECTS OF ETHEREAL ADDITIVES ON ANODIC FLUORINATION

In neat ionic liquid fluoride salts, the nucleophilicity of fluoride ions is rather low, resulting in poor fluorination yields. Fuchigami et  al. have successfully carried out the solvent-free electrochemical difluorodesulfurization of O-ethyl benzothioate in the presence of ether-containing additives like poly(ethylene glycol) (PEG) [94,95]. As shown in Equation 20.52, the addition of ca. 3% PEG additives to the reaction system greatly improved the yield due to its anodic stability and ability to coordinate the counter cations of fluoride ions. S

F

F

Additive

Yield (%)

DME PEG [Mn ~ 200]

22 18 80

–2e O

O Et4NF-3HF

© 2016 by Taylor & Francis Group, LLC

(20.52)

821

Fluorination

IV. ANODIC FLUORINATION OF MACROMOLECULES Electrochemical fluorination of conducting polymers has been demonstrated using anodic fluorodesulfurization of poly(fluorene) derivatives in ionic liquids Et4NF-nHF [96–98]. An alternating copolymer of 9,9-dioctylfluorene and 9,9-disulfanylfluorene gave a tough film on a platinum plate electrode by drop casting. This electrode was then placed in an undivided plastic cell filled with Et4NF-nHF, and constant current electrolysis was performed to yield poly(9,9-difluorofluorenealt-9,9-dioctylfluorene). The progress of the fluorodesulfurization reaction could be followed by NMR. The degree of fluorination gradually increased with the amount of charge passed and was mostly completed after passage of 24 F. The electrochemical reaction proceeded selectively even after a large amount of charge was passed (Equation 20.53) [96,97]. When the reaction was carried out in MeCN containing Et4NF-nHF as a supporting electrolyte, the polymer film detached from the platinum anode and is no longer electroactive. Thus, the choice of electrolytic media was highly important.

F

S

S

F

–2e, –2SAr

Octyl

Octyl

n

Et4NF-5HF 24 F

Octyl

Octyl

n

(20.53) Anodic fluorodesulfurization of the alternating copolymer of N-decylcarbazole and 9,9-disulfanylfluorene was investigated, and yielded poly(9,9-difluorofluorene-alt-N-decylcarbazole); however, the reaction resulted in only 50% conversion (Equation 20.54). In this case, the localization of cation charge on the carbazole unit probably prevented the discharge of the sulfanyl group [98].

S

S

C10H21 N –2e, –2SAr n

F

F

C10H21 N

(20.54)

Et4NF-5HF 24 F

n Conversion: 50%

V. INDIRECT ANODIC FLUORINATION As mentioned in Sections I, II.G.2 and II.G.3, anode passivation takes place very often, which results in poor yield and low current efficiency. In order to avoid such passivation, Fuchigami et al. developed an indirect electrochemical method using various mediators. Thus, Br+/Br− and triarylamine redox mediators have been shown to be effective for selective mono- and difluorodesulfurization of dithoacetals, respectively [99,100]. Furthermore, triarylamine has been shown to be a highly effective mediator for monofluorodesulfurization of β-lactams (Equation 20.55) [101]. In the absence of triarylamine, severe passivation of the anode takes place during anodic fluorination.

© 2016 by Taylor & Francis Group, LLC

822

Organic Electrochemistry Et3N-3HF

Anode

2F–

R1

SPh

+

2Ar3N

2e

N O

R2

R1

F

2Ar3N

N O

(20.55)

+ PhSF R2

Ar = 2,4-Br2C6H3 R1 = H, R2 = C6H5CH2: 83% R1 = H, R2 = p-BrC6H4CH2: 100% R1 = Me2SiOCH–, R2 = C6H5CH2: 66% t-Bu Me

They have also demonstrated the first catalytic use of hypervalent difluoroiodoarene anodically generated in situ for gem-difluorodesulfurization of dithioacetals, as shown in Equation 20.56 [102,103]. Pt anode

S

F MeO

2e

S

1/2

I F

X

2F–

X

F

(20.56)

F

I

MeO

1/2 1.9 V vs. SSCE

X

X X = Cl: 98% F: 96%

Even in ionic liquid HF salts, severe passivation of the anode often occurs. Therefore, a novel indirect electrochemical fluorination system was developed employing a task-specific ionic liquid with an iodoarene moiety as the mediator in HF salts [104]. The mediator improved the reaction efficiency for a variety of electrochemical fluorinations (Equation 20.57) and remained intact in the ionic liquids after the extraction process for reuse in subsequent runs. Et3N-3HF Mediator (10 mol%)

N N

Mediator: l

S

EWG

O

Tf2N– N N+

Undivided cell 4F 5 mA/cm2

N N

F S

EWG

(20.57)

EWG = COOEt, 87% (without mediator: 31%) EWG = CN, 72%

Task-specific ionic liquid with triarylamine mediator has been also developed for anodic fluorodesulfurization [105]. In addition, a polymer-supported iodobenzene (PSIB) mediator was also effective for indirect anodic fluorination in HF salts [106]. In this case, the iodobenzene moiety pendent from the solid polymer support could not be directly oxidized; therefore, a double mediator system was necessary. As shown in Equation 20.58, electrooxidation of Cl− gave Cl+, which reacted with the iodobenzene moiety to form PhI+Cl. This species then captured a fluoride ion to give the hypervalent [(chloro) (fluoro)iodo]benzene moiety. The hypervalent iodine moiety thus generated oxidizes the substrate,

© 2016 by Taylor & Francis Group, LLC

823

Fluorination

and consequently, the starting PSIB was recovered. The recovered PSIB mediator was reused in subsequent runs, maintaining a good yield (86–79%) of fluorinated product through to the 10th run. 1/2 F

F

Ph

Anode

1/2 S

Ph

S

Ph

Ph

Cl– F Ph

l

Ph

l Cl

(20.58)

–2e +

Cl

Ph

F–

l+ Cl

Ph

l

n

=

m

PSIB mediator l

VI.

ANODIC FLUORINATION USINg ALKALI METAL SALTS

Tajima et al. developed anodic fluorination based on cation exchange between alkali-metal fluorides and solid-supported acids like Amberlyst 15Dry [107]. However, in this system, anodically stable bases such as 2,6-lutidine must be added in order to increase the nucleophilicity of fluoride ions generated in situ. It is known that poly(ethylene glycols) [HO(CH2CH2O)nH, n > 3] are highly active and selective in catalyzing dehydrohalogenation in organic–aqueous hydroxide two-phase systems. Moreover, recently, it was demonstrated that tetraethylene glycol (terminal group: OH) could dissociate MF (M = K, Cs) into the fluoride ion and the metal cation in aprotic polar solvents. Based on these facts, Fuchigami et al. developed a novel electrolytic system and have achieved anodic fluorination of various organic compounds in tetraethylene glycol using alkali metal salts like KF (Equation 20.59) [108]. H

F

3 SPh 2 Z 3

Z = P, Sb

SPh

–ne

KF + PEG [Mn ~ 200] MeCN, Pt–Pt, r.t. ~40°C 4 F, 5 mA/cm2 Undivided cell

3 F

F

(20.59)

2 SPh

Z 3 F

VII. CONCLUSION In contrast to conventional chemical fluorination methods using hazardous and/or costly reagents such as F2, FClO3, CF3OF, XeF2, Et2NSF3 (DAST), N-fluoropyridinium salts, N-fluorotriethylenediamine derivative (Selectfluor®), and 4-tert-butyl-2,6-dimethylphenylsulfur trifluoride (Fluolead™), selective anodic fluorination can be readily carried out under mild conditions and does not require any hazardous reagents. As described earlier, selective electrochemical fluorination is a promising procedure for the preparation of various partially fluorinated organic substances.

© 2016 by Taylor & Francis Group, LLC

824

Organic Electrochemistry

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Fluorination 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. 61. 62. 63. 64. 65. 66. 67. 68. 69. 70. 71. 72. 73. 74. 75. 76. 77. 78. 79. 80. 81. 82. 83. 84. 85. 86. 87. 88. 89. 90. 91. 92. 93. 94. 95. 96. 97. 98. 99. 100. 101. 102. 103.

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104. Sawamura T., Kuribayashi S., Inagi S., Fuchigami T. Org. Lett. 2010, 12, 644–646. 105. Takahashi K., Furusawa T., Sawamura T., Kuribayashi S., Inagi S., Fuchigami T. Electrochim. Acta 2012, 77, 47–53. 106. Sawamura T., Kuribayashi S., Inagi S., Fuchigami T. Adv. Synth. Catal. 2010, 352, 2757–2760. 107. Tajima T., Nakajima A., Doi Y., Fuchigami T. Angew. Chem., Int. Ed. 2007, 46, 3550–3552. 108. Sawamura T., Takahashi K., Inagi S., Fuchigami T. Angew. Chem., Int. Ed. 2012, 51, 4413–4416.

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Section V Electrochemical Conversions of Organic Compounds

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21

Electrochemistry of Fullerenes, Derivatives, and Related Compounds Frederic Melin, Lourdes E. Echegoyen, and Luis Echegoyen

CONTENTS I. Introduction ............................................................................................................................. 830 II. General Trends in the Redox Properties of Fullerenes and Related Compounds .................. 831 A. Cage Effects in Empty Fullerenes .................................................................................. 831 B. Cage and Encapsulated Species Effects in Endohedral Metallofullerenes .................... 831 1. Monometallofullerenes ............................................................................................ 832 2. Dimetallofullerenes.................................................................................................. 833 3. Clusterfullerenes ...................................................................................................... 835 C. Influence of Exohedral Derivatization ............................................................................ 839 1. Empty Fullerenes ..................................................................................................... 839 2. Endohedral Metallofullerenes Sharing the C80 (Ih) Cage ........................................840 3. M@C82 (C2v) Endohedral Metallofullerenes ............................................................ 841 4. Endohedral Metallofullerenes Incorporated in Donor–Acceptor Dyads................. 843 D. Larger Carbon Nanostructures: Structure, Defects, and Impurity Effects ..................... 845 III. Preparation, Purification, and Derivatization of Fullerenes and Related Compounds Based on Their Redox Properties ...........................................................................................846 A. Isolation of Kinetically Unstable Empty and Endohedral Fullerenes ............................846 1. Looking for the Missing Empty Carbon Cages .......................................................846 2. Separating Endohedral Metallofullerenes from Empty Fullerenes ......................... 847 3. Looking for the Missing Endohedral Metallofullerenes .........................................848 B. Separation of Isomeric Carbon Cages ............................................................................848 1. By Selective Oxidation.............................................................................................848 2. Using the Retrocyclopropanation Reaction.............................................................. 849 C. Preparation of Exohedral Derivatives of Fullerenes ....................................................... 849 1. By Nucleophilic Substitution ................................................................................... 849 2. Using the Retrocyclopropanation Reaction.............................................................. 850 3. Fine-Tuning the Reactivity of Endohedral Fullerenes ............................................. 850 D. Solution Processing of Larger Carbon Nanostructures .................................................. 852 IV. Conclusion............................................................................................................................... 852 References ...................................................................................................................................... 853

829

© 2016 by Taylor & Francis Group, LLC

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I. INTRODUCTION The most intriguing property of buckminsterfullerene C60, discovered in 1985, is without doubt its outstanding ability to accept electrons with very low reorganization energy. Using either cyclic voltammetry (CV), differential pulse voltammetry (DPV), or Osteryoung square wave voltammetry (OSWV), six reversible waves are observed for C60 at low temperature (Figure 21.1) [1], which was predicted by theory, based on its energetically low-lying triply degenerate lowest unoccupied molecular orbital (LUMO) [2]. Each one of the reduction steps has a relatively constant separation between them of 450 ± 50 mV. Most importantly, with each addition of one electron, no Jahn– Teller distortion is observed, that is, the C60 cage does not suffer any structural modification from its highly symmetrical, quasi-spherical architecture [3]. This latter characteristic is necessary for the long-lived charge-separated states required for photovoltaic applications [4], an area that has recently made the fullerene field one of the most active in chemistry. The sister molecule, C70, exhibits a very comparable electron-accepting behavior, except that all six waves are observable at room temperature [1]. The larger fullerenes also possess very rich electrochemistry and high symmetry and exhibit no Jahn–Teller distortion. Consequently, they have also attracted considerable interest in the scientific community in the last two decades. Numerous studies comprising members of the fullerene family are focusing on tuning their redox chemistry by surface functionalization or by trapping metals and clusters inside the cages. The search for compounds with encaged moieties (the so-called endohedral fullerenes) has resulted in the synthesis of carbon cages with unusual sizes and symmetries, including some with fused pentagon rings [5]. Fullerene molecules with carefully designed physicochemical properties have been produced by functionalization of various carbon cages [6]. These new molecules all possess redox properties that can be used either as a synthetic tool to prepare, separate, or purify them or as an analytical tool to characterize and classify the carbon cages, predict their reactivity, or understand the electronic interplay between encapsulated species and carbon cages in the case of endohedral fullerenes. In this chapter, we will first review the main trends in the electrochemical properties of fullerenes and related compounds, with a special emphasis on the growing family of endohedral fullerenes. We will then focus on representative examples that show how these unique redox properties can C60 at –10°C 10 μA

(a)

5 μA

–1.0 (b)

–2.0 Potential (V vs. Fc/Fc+)

–3.0

FIgURE 21.1 CV (a) and OSWV (b) of C60 in acetonitrile/toluene + 0.1 M (n-Bu)4NPF6 at −10°C using glassy carbon working electrode (GCE) and ferrocene/ferrocenium (Fc/Fc+) couple as internal reference. (Reprinted from Xie, Q. et al., J. Am. Chem. Soc., 114, 3978, 1992. With permission.)

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be exploited to prepare some of these compounds. Due to space limitations, it is impossible to be exhaustive on this rapidly expanding research domain. Therefore, we suggest the reader to also consult several comprehensive reviews for more details on this subject [7–8].

II.

gENERAL TRENDS IN THE REDOX PROPERTIES OF FULLERENES AND RELATED COMPOUNDS

A.

CAGE EFFECTS IN EMPTY FULLERENES

The electrochemical properties of fullerenes are highly dependent on the cage size and symmetry. Early studies dealt mainly with empty carbon cages under various solvents, supporting electrolytes and temperature conditions. A noteworthy example is the comparison between C60 and C70. Their corresponding first and second reduction potential values are nearly identical in acetonitrile/toluene. However, from the trianion to the hexa-anion, C70 becomes increasingly easier to reduce than C60. In other solvents, the differences, although still small, may be slightly more pronounced (see Table 21.1). Of electrosynthetic relevance is the fact that C60 anions are stable on the voltammetric time scale, but when electrogenerated by controlled potential coulometry in solution, only the mono- through tetra-anions are stable [7a,b]. Furthermore, neutral C60 is insoluble in solvents such as DMF, acetonitrile, or THF, but its anions dissolve readily. Therefore, these anions can be generated from a suspension of C60 [9]. These properties have been used to prepare several C60 derivatives starting from the electrogenerated anions [11]. The soot collected from graphite arcing yields, albeit in very small quantities after separation, isomerically pure samples of C76, C78, C82, and C84 and isomeric mixtures of C86, C90, C92, and C96. Studies on the cathodic electrochemistry of C60 to C84 indicate a very distinct trend: as the size of the cage increases, the first reduction is systematically shifted toward more positive potentials. However, there is no apparent trend for the first oxidation potential (see Table 21.2). Generally, the larger cages have a lower electrochemical gap ΔEgap than C60 and C70. Beyond C84, unfortunately, only electrochemical studies of isomeric mixtures are available. As an example, C92, as isolated and purified by high-performance liquid chromatography (HPLC), exhibits eight reversible reductions that can be grouped into two distinct sets, based on their intensity and assigned to two different cage isomers [12]. This study clearly demonstrates that the electrochemical properties of two fullerene isomers can be dramatically different.

B.

CAGE AND ENCAPSULATED SPECIES EFFECTS IN ENDOHEDRAL METALLOFULLERENES

Most of the current reports on fullerene electrochemistry focus on endohedral metallofullerenes (EMFs) with the aim to elucidate the electronic interplay between encaged species and carbon cages. A substantial amount of knowledge has been obtained through systematic studies of simple metallofullerene families such as the M@C82 (M = La, Pr, Gd, Ce, Tm, and Y) [14,15], Yb [16], TAbLE 21.1 Half-Wave Reduction Potentials (vs. Fc/Fc+) of C60 and C70 in Acetonitrile/Toluene (PhMe/MeCN) and Dichloromethane (DCM) Compound C60 C70 C60 C70

Solvent

Temperature (°C)

E1

E2

E3

E4

E5

E6

Reference

PhMe/MeCN PhMe/MeCN DCM DCM

−10 −10 +25 +25

−0.98 −0.97 −1.02 −0.93

−1.37 −1.34 −1.41 −1.31

−1.87 −1.78 −1.87 −1.73

−2.35 −2.21

−2.85 −2.70

−3.26 −3.70

[1] [1] [9] [10]

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−2.09

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Organic Electrochemistry

TAbLE 21.2 Redox Potentials versus Fc/Fc+ of Some Isomerically Pure Empty Fullerenes in 1,1,2,2-Tetrachloroethane Unless Otherwise Stated Compound C60 C70 C76 C78 (C2v) C82a C84 a b

E1ox (V )

E1red (V )

ΔEgap (V) b

Reference

1.26 1.20 0.81 0.95 0.72 0.93

−1.06 −1.02 −0.83 −0.77 −0.69 −0.67

2.32 2.22 1.64 1.72 1.41 1.60

[13] [13] [13] [13] [14a] [13]

1,2-Dichlorobenzene. Electrochemical bandgap ΔEgap = E1ox − E1red.

Sm [17], and Ca@C2n (n = 37–48) [18], as well as the trimetallic nitride cluster fullerenes M3N@C2n (M = Sc, Y, and most lanthanides, n = 34–48) [19]. 1. Monometallofullerenes In the simple metallofullerenes, there is a clear dependency of the electrochemical properties on the number of electrons transferred by the metal to the cage (i.e., the oxidation number of the metal). Sc, Y, and the majority of the lanthanides can give three electrons to the cage and form the so-called trivalent EMFs. It seems that these metals are preferentially encapsulated inside the C2v and Cs isomers of C82, with the C2v as the most abundant [20]. These metallofullerenes are both easier to oxidize and easier to reduce than the empty fullerenes and their first redox processes are reversible. Their strong electron-accepting behavior is quite unexpected, when considering that the C82 cage already bears a formal 3− charge. Their small electrochemical HOMO–LUMO gap (less than 0.5 V, see Table 21.3) is a remarkable characteristic and reflects their open-shell electronic structure. Overall, their oxidation state can be changed from 2+ down to 6−. Interestingly, no major differences exist in the electrochemical properties of the C2v and Cs isomers, which suggests that their behavior is mostly controlled by their radical nature and not by the structure of the cage. However, good linear relationships between the first redox potentials of these metallofullerenes and the ionic radii of the encapsulated metals were found (see Figure 21.2) [14a]. These relationships can be understood by taking into account the probable structure of these endohedral fullerenes: that is, a positively TAbLE 21.3 Half-Wave Potentials versus Fc/Fc+ of Trivalent Mono Endohedral Metallofullerenes Measured in 1,2-Dichlorobenzene Fullerene

ox2 E1/2

ox1 E1/2

red1 E1/2

red2 E1/2

red3 E1/2

red4 E1/2

ΔEgap a

Reference

La@C82 (C2v) La@C82 (Cs) Pr@C82 (C2v) Pr@C82 (Cs) Ce@C82 (C2v) Gd@C82 (C2v) Y@C82 (C2v)

1.07 1.08 1.08 1.05 1.08 1.08 1.07

0.07 −0.07 0.07 −0.07 0.08 0.09 0.10

−0.42 −0.47 −0.39 −0.48 −0.41 −0.39 −0.37

−1.37 −1.40 −1.35 −1.39 −1.41 −1.38 −1.34

−1.53 −2.01 −1.46 −1.99 −1.53 −2.22 −2.22

−2.26 −2.40 −2.21 −1.99 −1.79

0.49 0.54 0.46 0.55 0.49 0.48 0.47

[14a] [15b] [15b] [15b] [14a] [14a] [14a]

a

1 ΔEgap = E1ox/ 21 − E1red /2 .

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−2.47

Electrochemistry of Fullerenes, Derivatives, and Related Compounds

0.10

Y

E ox1

Y

E red1

833

Gd Ce

E (V) vs. Fc/Fc+

0.08

0.06

La

–0.38 Gd –0.40 Ce –0.42 La –0.44

1.16

1.12 1.08 Ionic radius (Å)

1.04

FIgURE 21.2 Linear relationships between the redox potential of the La@C82 C2v compounds and the ionic radii of M3+. (Reprinted from Suzuki, T. et al., Tetrahedron, 52, 4973, 1996. With permission.)

charged metal (3+) that is not centered on but which strongly interacts with a hexagonal ring of the cage along the C2 axis. The smaller the metal, the closer to the cage it can be and the tighter the electrons are bound to the cage. Small metals such as Y thus form EMFs that are easy to reduce and difficult to oxidize, whereas larger metals such as the lanthanides form EMFs that are easier to oxidize but more difficult to reduce. In contrast, alkaline earth metals such as Ca as well as some lanthanides such as Sm, Yb, and Eu transfer only two electrons to the cage, and therefore form divalent EMFs. Their electron-accepting character (see Table 21.4) is also stronger than that of the empty fullerenes. The large gap between the second and third reduction steps is indicative of nondegenerate LUMO-1 and LUMO-2 orbitals [16a]. Most of the electrochemical studies of these compounds were done in a mixture of toluene and acetonitrile, and no oxidation was observed in this solvent system. Their electrochemical gaps were therefore predicted to be larger than those of the trivalent EMFs, which was consistent with their believed closed-shell electronic structure. In 1,2-dichlorobenzene, however, one or two quasi-reversible oxidation waves were recently observed for most ytterbium- and a few samarium-based compounds. These studies led to the determination of the electrochemical bandgaps of these metallofullerenes (see Table 21.4) and to a correlation between these values and the production yields [16b]. Divalent EMFs are found with quite a large diversity of cage sizes and symmetries, and the influence of the cage structure on the electrochemical properties is very significant in this family of compounds. As shown in Table 21.4, dramatic differences in the electron-accepting abilities of isomeric endohedral fullerenes are sometimes observed. For instance, the C2 isomer of Sm or Yb@C82 is more difficult to reduce than the Cs or C2v isomers by at least 0.25 V! When the cage becomes larger, the number of oxidation and reduction processes also tends to increase. In contrast, divalent EMFs with the same carbon cage but different metals usually exhibit very close electron-accepting abilities, which suggests that the metal hardly contributes to the LUMO of these compounds. 2. Dimetallofullerenes Dimetallofullerenes mainly with trivalent metals such as Sc [21], La [22], Er [23], Ce [22e,24], and cages as small as C72 and as large as C82 have been characterized. In these structures, it was presumed

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Organic Electrochemistry

TAbLE 21.4 Half-Wave Potentials against Fc/Fc+ of Divalent Mono Endohedral Metallofullerenes Measured in Toluene/Acetonitrile Unless Otherwise Stated Fullerene

ox2 E1/2

ox1 E1/2

red1 E1/2

red2 E1/2

red3 E1/2

red4 E1/2

ΔEgapb

Reference

Sm@C74 (D3h) Sm@C76 (I) Sm@ C80 (C2v(3))

— — 0.85

— — 0.43a

— — — — — — — — — — — — — — — — 0.78a

— 0.42a

— — 0.90a

0.34a

0.61a 0.46a

Yb@C84 (C1(12))

— — — — 0.68a

Yb@C84 (C2(11))

0.48a

0.19a

Ca@C76 Ca@C82 (C2) Ca@C82 (C3v(7)) Ca@C84 (C2(13))

— — — —

— — — —

−0.98 −0.85 −0.98 −1.23a −0.59 −0.76 −1.01a −0.57 −0.87 −0.70 −0.78 −0.87 −0.86 −0.61 −0.84 −0.85 −0.96 −0.83 −1.02 −0.79 −0.95 −1.27a −0.65 −0.92a −0.76 −0.98a −0.67 −0.78a −0.88 −1.16a −0.68 −0.94a −0.72 −1.14a −0.99 −0.74 −0.96 −0.90

−1.55 −1.43 −1.58 −1.76a −1.54 −1.27 −1.51a −1.45 −1.12 −1.58 −1.31 −1.42 −1.30 −1.11 −1.33 −1.29 −1.55 −1.46 −1.59 −1.46 −1.55 −1.87a −1.58 −1.81a −1.33 −1.50a −1.56 −1.60a −1.26 −1.50a −1.57 −1.76a −1.34 −1.70a −1.57 −1.30 −1.55 −1.27

−1.96 −1.74 −1.92 −2.07a −1.80 −1.66 −1.90a −1.81 −1.31 −1.81 −1.71 −1.78 −1.65 −1.28 −1.69 −1.50 −1.99 −1.89 −2.01 −1.83 −1.90 −2.13a −1.81 −2.01a −1.73 −1.87a −1.90 −1.90a −1.64 −1.86a −1.79 −1.97a −1.54 −2.06a −1.97 −1.70 −1.90 −1.65

— — 1.28a

Sm@C82 (Cs(6)) Sm@ C82 (C2(5))

−0.52 −0.45 −0.59 −0.85a −0.27 −0.54 −0.84a −0.22 −0.66 −0.39 −0.47 −0.62 −0.54 −0.40 −0.51 −0.52 −0.52 −0.46 −0.68 −0.48 −0.57 −0.89a −0.33 −0.62a −0.60 −0.86a −0.33 −0.46a −0.63 −0.95a −0.49 −0.76a −0.46 −0.85a −0.61 −0.59 −0.65 −0.64

[17a] [17a] [17a] [17c] [17a] [17a] [17b] [17a] [17a] [17a] [17a] [17a] [17a] [17a] [17a] [17a] [16a] [16a] [16a] [16a] [16a] [16b] [16a] [16b] [16a] [16b] [16a] [16b] [16a] [16b] [16a] [16b] [16a] [16b] [18] [18] [18] [18]

Sm@C82 (C2v(9)) Sm@C84 (C2(13)) Sm@C84 (C1(12)) Sm@C86 Sm@C90 (C2(40)) Sm@C90 (C2(42)) Sm@C92 (C1(42)) Sm@C94 (C3v(134)) Sm@C96 Yb@C74 (D3h) Yb@C76 (I) Yb@C76 (II) Yb@C78 Yb@C80 (C2v(3)) Yb@C82 (Cs(6)) Yb@C82 (C2(5)) Yb@C82 (C2v(9)) Yb@C84 (C2(13))

a b

— — — — — — — — — — — — — 0.34a

0.38a

0.22a

— 1.26a — — — — — — — — — — — — — 1.23a 0.96a 1.24a 1.07a 1.41a 0.98a 1.04a — — — —

In 1,2-dichlorobenzene. 1 ΔEgap = E1ox/ 21 − E1red / 2 . Attribution of the carbon cage symmetries is mostly based on recent x-ray studies [20].

that the carbon cage accepts six electrons from the encapsulated atoms and should therefore not be prone to further reduction. Unexpectedly, these endohedrals exhibit both strong electron-accepting and electron-donating abilities and thus relatively low electrochemical bandgaps (see Table 21.5). In particular, La2@C80 and Ce2@C80 are even easier to reduce than La@C82 and Ce@C82 as a result of the LUMO being localized primarily on the encapsulated atoms. The M2@C72 compounds are

© 2016 by Taylor & Francis Group, LLC

835

Electrochemistry of Fullerenes, Derivatives, and Related Compounds

TAbLE 21.5 Half-Wave Potentials against Fc/Fc+ of Dimetallo-Endohedral Fullerenes Measured in 1,2-Dichlorobenzene + 0.05 M (n-bu)4NPF6 Unless Otherwise Stated Compound

Eox

La2@C72 (D2) Ce2@C72 (D2) La2@C78 (D3h) La2@C80 (Ih) La2@C80 (D5h) Ce2@C80 (Ih) Ce2@C80 (D5h) Sc2@C82 (C3v (8)) Er2@C82b

0.75 0.82 0.62 0.95 0.78 0.95 0.66 — —

a b

2

Eox

1

0.24 0.18 0.26 0.56 0.22 0.57 0.20 0.05 0.19

Ered

Ered

Ered

ΔEgapa

Reference

−0.68 −0.81 −0.40 −0.31 −0.36 −0.39 −0.40 −1.10 −0.87

−1.92 −1.86 −1.84 −1.72 −1.72 −1.71 −1.76 — −1.26

— — −2.28 −2.13 — — −2.16 — —

0.92 0.99 0.66 0.87 0.58 0.96 0.60 1.15 1.06

[22a] [24a] [22b] [22c,d] [22e] [24b] [22e] [21] [26]

1

2

3

1 ΔEgap = E1ox/ 21 − E1red /2 . In pyridine.

among the few examples of kinetically stable EMFs with a cage that violates the Isolated Pentagon Rule (non IPR fullerenes) [25]. The stabilization of the fused pentagon system in these structures is believed to occur through specific coordination with the metallic moiety and some degree of intramolecular ion pairing [22a]. 3. Clusterfullerenes Among the EMFs, the trimetallic nitride cluster fullerenes M3N@C2n (MNEFs) currently constitute the most abundant and diverse, probably thanks to their larger electrochemical bandgaps, which range from 1.10 to 2.08 V (see Table 21.6) [19]. Indeed, most of the lanthanides, except Sm, Yb, and Eu, as well as Sc and Y form such compounds. After C60 and C70, Sc3N@C80 is the third most abundant fullerene that can be prepared in an arc reactor, which was totally unexpected when considering that on their own, the C80 cage and Sc3N cluster are highly unstable [27]. In these endohedrals, it is widely accepted that stability is reached through a transfer of six electrons from the cluster to the cage, which requires a metal with a formal oxidation number of 3+. It is not a coincidence that Sm, Yb, and Eu, which apparently do not form trimetallic nitride fullerenes, prefer also to form divalent instead of trivalent EMFs (see Section II.B.1). For lanthanides smaller than Gd [28], the most abundant cage formed is IPR isomer 7 of C80 with (Ih) symmetry, and thus a large number of M3N@C80 (Ih) are available (see Table 21.6). For larger metals such as Nd, Pr, and Ce [29], C88 (presumably IPR isomer 35 with D2 symmetry) is preferred, while La prefers C96 (possibly IPR isomer 186 with D2 symmetry) [30]. Several of these compounds feature a non IPR-cage, including Sc3N@C68 (D3) [31], Sc3N@C70 (C2v) [32], Gd3N@C78 (C2) [33], Gd3N@C82 (Cs) [34], and Gd3N@C84 (Cs) [35]. MNEFs usually exhibit one or two reversible oxidation steps and at least two irreversible reduction steps using CV, while their chemical reduction is reversible. The electrochemical irreversibility suggests that an internal structural reorganization occurs upon addition of electrons, which may hinder the free rotation of the cluster [36]. Only Sc3N@C80 at high scan rates [36a], TiSc2N@ C80 [37a], TiSc2N@C80 [37b], and the C88 MNEFs [29,38] (see Figure 21.3) show reversible behavior for both oxidation and reduction scans. The nature of the metal does not appear to influence significantly the oxidation and reduction potentials, and hence, it does not affect the value of the electrochemical gap of the MNEFs (see Figure 21.3 and Table 21.6). These observations suggest that both the HOMO and the LUMO of MNEFs are mainly cage localized orbitals. Only in the case of Sc3N@C80, it was demonstrated that the  cluster contributes significantly to the LUMO [39]. Mixed metal nitride compounds based on

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836

Organic Electrochemistry

TAbLE 21.6 Reductive Cathodic Peak Potentials (Epc) and Oxidative Half-Wave Potentials (E1/2) against Fc/Fc+ of Trimetallic Nitride Endohedral Fullerenes as Obtained by Cyclic Voltammetry (Unless Otherwise Stated) in 1,2-Dichlorobenzene + 0.05 M (n-bu)4NPF6 Compound Sc3N@C68 (D3) Sc3N@C78 (D3h) Y3N@C78 (C2) Dy3N@C78 (C2) Gd3N@C78 (C2) Sc3N@C80 (Ih) Y3N@C80 (Ih) Lu3N@C80 (Ih) Tm3N@C80 (Ih) Er3N@C80 (Ih) Ho3N@C80 (Ih) Dy3N@C80 (Ih) Tb3N@C80 (Ih) Gd3N@C80 (Ih) Nd3N@C80 (Ih) Pr3N@C80 (Ih) ScYErN@C80 (Ih) CeSc2N@C80 (Ih) CeY2N@C80 (Ih) CeLu2N@C80 (Ih) PrSc2N@C80 (Ih) TiSc2N@C80 (Ih) TiY2N@C80 (Ih) Sc3N@C80 (D5h) Lu3N@C80 (D5h) Tm3N@C80 (D5h) Dy3N@C80 (D5h) Gd3N@C82 (Cs) Gd3N@C84 (Cs) Nd3N@C84 Gd3N@C86 (D3) Nd3N@C86 Pr3N@C86 Lu3N@C88 Gd3N@C88 Nd3N@C88 Pr3N@C88 Ce3N@C88 La3N@C88 Pr3N@C92 Ce3N@C92 La3N@C92 Pr3N@C96 Ce3N@C96 La3N@C96 a b

ox2 E1/2

0.85 0.68 0.53 — — 1.09 — 1.11 1.15 — — — — — — — — — — — — — — — — 0.79 — — — — — — — 0.44 0.49 0.53 0.54 0.63 0.66 0.79b 0.76b — 0.53 0.67 0.53

1 . ΔEgap = E1ox/ 21 − E1ox/ 21 − Epred c OSWV peaks.

© 2016 by Taylor & Francis Group, LLC

ox1 E1/2

0.33 0.21 0.25 0.47 0.47 0.59 0.64 0.64 0.65 0.63 0.60b 0.70 0.59b 0.58 0.63 0.59 0.64 0.33 −0.07 0.01 0.64 0.16 0.00 0.34 0.45 0.39 0.40 0.37 0.32 0.31 0.35 0.36 0.31 0.02 0.06 0.07 0.09 0.08 0.21 0.35 0.35 0.36 0.14 0.18 0.14

1 Epred c

2 Epred c

3 Epred c

ΔEgapa

Reference

−1.45 −1.56 −1.62 −1.54 −1.53 −1.26 −1.41 −1.42 −1.43 −1.40 −1.45b −1.37 −1.38b −1.44 −1.42 −1.41 −1.55 −1.34 −1.36 −1.43 −1.32 −0.94 −1.11 −1.33 −1.41 −1.45 −1.40 −1.52 −1.37 −1.44 −1.35 −1.46 −1.48 −1.35 −1.43 −1.36 −1.34 −1.33 −1.36 −1.46 −1.46 −1.44 −1.51 −1.50 −1.54

−2.05 −1.91 −1.99 −1.93 −1.89 −1.62 −1.83 −1.80 −1.78 −1.83 −1.91b −1.86 −1.86b −1.86 −1.89 −1.84 −1.97 −1.87 −1.88 −1.92 −1.91 −1.58 −1.79 — — −1.81 −1.85 −1.86 −1.76 −2.02 −1.70 −1.79 −1.80 −1.68 −1.74 −1.75 −1.72 −1.60 −1.67 −1.82 −1.82 −1.69 −1.86 −1.84 −1.77

— — — — — −2.37 −2.34 −2.26 — −2.16 −2.21b — −2.16b −2.18 — — — — — — — −2.21 — — — −2.36 — — — — — — — — — −2.39 −2.35 −2.35 — — −2.36 −2.29 — — —

1.78 1.77 1.87 2.01 2.00 1.85 2.05 2.06 2.08 2.05 2.05 2.07 1.97 2.02 2.05 2.00 2.19 1.67 1.30 1.44 1.96 1.10 1.11 1.67 1.86 1.84 1.80 1.89 1.69 1.75 1.70 1.72 1.79 1.37 1.49 1.43 1.43 1.41 1.57 1.81 1.80 1.80 1.65 1.68 1.68

[42] [33] [43] [36b] [33] [36a] [44] [45] [46] [44] [47] [36b] [47] [38] [29b] [29b] [48] [40b] [40b] [40a] [40b] [37a] [37b] [49] [49] [46] [36b] [50] [50] [29b] [50] [29b] [29b] [51] [50] [29a] [29b] [29b] [30] [41b] [41b] [41b] [41b] [41b] [41b]

837

Electrochemistry of Fullerenes, Derivatives, and Related Compounds

Ce3N@C88

Current (arbitrary unit)

Pr3N@C88

Nd3N@C88

Gd3N@C88

1.0

FIgURE 21.3 rate 0.1 Vs−1).

0.5

0.0

–0.5 –1.0 Potential (V vs. Fc+/Fc)

–1.5

–2.0

–2.5

CV of M3N@C88 (M = Gd, Nd, Pr, and Ce) obtained in o-DCB + 0.05 M (n-Bu)4 NPF6 (scan

Ce and Ti metals constitute other exceptions to this rule. For TiM2N@C80 compounds, the valence state of the Ti atom is changed during both oxidation and reduction processes [37], whereas for CeM2N@C80 compounds [40], oxidation occurs at the cerium metal and the corresponding redox potential is highly dependent on the size of the second metal. Table 21.6 also emphasizes that the first oxidation potential is very sensitive to the size and symmetry of the cage (see in particular the difference between the (Ih) and (D5h) isomers of C80 MNEFs), whereas the first reduction potential is only slightly influenced by the nature of the cage. Poblet et al. computed the orbital energies of the cages by the density functional theory (DFT) and uncovered an excellent correlation between the electrochemical gaps of the MNEFs and the (LUMO-4)–(LUMO-3) orbital gaps of the free cages (Figure 21.4), which nicely confirmed the ionic bonding model M3N6+@C806− proposed for these compounds [39,41]. Indeed, with such an electron transfer, the HOMO and the LUMO of the MNEF should correspond respectively to the LUMO-3 and the LUMO-4 of the cage. In addition, the cluster seems to select the cages with the largest (LUMO-4)–(LUMO-3) orbital gaps. Other reported encapsulated clusters inside fullerene cages include metal carbides (M2C2, M3C2, and M4C2; M = Sc, Ti, Y, Lu, Dy, and Gd) [52] and more recently metal oxides (M4O2 and M4O3, M  =  Sc) [53], metal sulfides (M2S, M = Sc, Ti, Y, Dy, and Lu) [54], as well as unique Sc3NC [55], Sc3CH [56], and YCN [57] moieties. The remarkable feature of the metal carbide endohedral fullerenes is the high flexibility of charge displayed by the C2 moiety. The electronic structures of Sc2C2@C82, Sc3C2@C80, and Sc4C2@C80 have been shown by DFT calculations to be, respectively, (Sc3+)2C22− C824−[58], (Sc3+)3C23− C806− [59], and C26−(Sc3+)4C806− [59,60]. The unusual C23− moiety confirms the ability of carbon cages to accommodate in their interior some species that cannot be isolated in regular chemical environments. More surprisingly, it seems that reduction or oxidation of Sc3C2@C80 affects only the charge of the C2 moiety, while the Sc and the cage keep their respective charges of 3+ and 6− [59]. As usual, a low electrochemical bandgap (0.50 V) characterizes

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838

Electrochemical gaps M3N@C2n/V

Organic Electrochemistry

C80

R2 = 0.98

2

C92

Sc

1.8 C96

C86

Sc Y Ce Pr Nd Gd Lu La Lineal (average)

C84

1.6

1.4 C88 1.0

0.8

1.2 1.4 1.6 Orbital gap free cage/eV

1.8

2.0

FIgURE 21.4 Linear relationship between the electrochemical gaps of MNEFs and calculated (LUMO-4)– (LUMO-3) orbital gaps of the corresponding empty fullerenes. (Reprinted from Chaur, M.N. et al., Angew. Chem. Int. Ed., 48, 1425, 2009. With permission.)

species with an open-shell electronic configuration (such as Sc3C2@C80 (Ih)). Compounds with the same carbon cage usually exhibit striking similarities in their redox behavior provided their clusters are valence isoelectronic. The analogous electrochemical properties of Sc2C2@C82, Sc2O@C82, and Sc2S@C82 (Cs(6)), for instance, can be understood on the basis of a donation of 4 electrons to the cage by the cluster in each case (see Table 21.7). Conversely, the electrochemical properties of Sc3C2@C80 (Ih) and Sc4C2@C80 (Ih) are distinct despite the fact that they share the same carbon cage. The differences arise from the Sc3C2 and Sc4C2 clusters that could have an open-shell [59] and a closed-shell configuration, respectively [59,60]. Similarly, the electrochemical behavior of Sc2C2@ C82 (C3v(8)) is very different from that of Sc2@C82 (C3v(8)) [21]. Overall these new clusterfullerenes TAbLE 21.7 Redox Potentials against Fc/Fc+ of Metal Carbide, Metal Oxide, Metal Sulfide Cluster Fullerenes, and Related Compounds Measured in 1,2-Dichlorobenzene Fullerene

Eox

Eox

Sc2C2@C82 (C3v (8)) Sc2C2@C82 (Cs (6)) Sc2C2@C72(Cs (10528)) Sc3C2@C80 (Ih) Sc4C2@C80 (Ih) Lu3C2@ C88 (D2) Sc2O@C82 (Cs (6)) Sc4O2@C80 (Ih) Sc2S@C70 (C2(7892)) Sc2S@C72 (Cs (10528)) Ti2S@C78 (D3h (10528)) Sc2S@C82 (Cs (6)) Sc2S@C82 (C3v (8)) Sc3NC@C80 (Ih (7)) Sc3NC@C78 (C2) YCN@C82 (Cs(6))

— 0.64 — — 1.10 — 0.72

0.47 0.42 0.41 −0.06 0.40 0.31 0.35 0.00 0.14 0.64 0.23 0.39 0.52 0.60 0.57 0.56

© 2016 by Taylor & Francis Group, LLC

2

0.65 0.96 0.65 0.65 0.96 — —

1

Ered

1

−0.94 −0.93 −1.18 −0.50 −1.16 −1.34 −0.96 −1.10 −1.44 −1.14 −0.92 −0.98 −1.04 −1.05 −1.05 −0.59

Ered

Ered

ΔEgap

Reference

−1.15 −1.30 −1.54 −1.64 −1.65 −1.70 −1.28 −1.73 −1.87 −1.53 −1.53 −1.12 −1.19 −1.68 −1.81 −0.84

−1.60 — −1.75 −1.82 — −2.15 −1.74 −2.35 −1.99 −2.24 −1.80 −1.73 −1.63 — — −1.76

1.41 1.35 1.59 0.44 1.56 1.65 1.31 1.10 1.58 1.78 1.15 1.37 1.56 1.65 1.62 1.15

[58a] [61] [62] [63] [60] [64] [54c] [53c] [54e] [54d] [54f] [54c] [54c] [55a] [55b] [57]

2

3

Electrochemistry of Fullerenes, Derivatives, and Related Compounds

839

exhibit smaller electrochemical bandgaps and thus lower thermal stability than the corresponding MNEFs with the same cage. This highlights once again the unique electronic interplay between the trimetallic nitride clusters and the carbon cages that results in the exceptional stability of the MNEFs.

C.

INFLUENCE OF EXOHEDRAL DERIVATIzATION

1. Empty Fullerenes The electrochemical behavior of C60 mono or multiadducts has been the subject of numerous studies and several reviews [7a–c,65]. We will thus recall here only the main conclusions of these studies for comparison purposes with recent reports for endohedral fullerene derivatives. Most C60 adducts and cycloadducts can be classified into either 1,2- or 1,4-derivatives if the new bonds are created respectively on adjacent carbons or on opposite carbons of a 6-membered ring. 1,2-Derivatives can be further divided into [6,6] or [5,6] adducts if the new bonds are formed respectively at the junction between two 6-membered rings or at the junction between a 5- and a 6-membered ring. The [5,6] adducts are usually fulleroids, because the junction needs to be broken to form the new bonds. Except in the case of [5,6] open adducts [66], derivatization results in partial loss of conjugation, from 60 to 58π electrons, which usually leads to a cathodic shift of both oxidation and reduction waves (between 30 and 350 mV per adduct) [7a,c]. However, the electronic properties of fullerenes can be finely tuned by using various types of addends. Those containing electron-accepting groups (such as −CN, −NO2, −F, ammonium salts) [67] or electronegative heteroatoms (O, N) [68] directly attached to the cage increase the electron affinity of C60, whereas those containing electropositive atoms (such as Si) [69] improve its electron-donating ability. For these derivatives, it is very important to carefully analyze any irreversible electrochemical features since these can be an indication that the adduct is not stable when submitted to a redox process. Echegoyen et al. established that some of the most commonly prepared C60 cycloadducts, that is, di(alkoxycarbonyl)-methano (also known as Bingel adducts) [70] and pyrrolidino [71] derivatives, can undergo a retrocycloaddition reaction under reductive and oxidative conditions, respectively. In CV studies, the methano adducts exhibit an irreversible second reduction process, whereas the pyrrolidino adducts exhibit an irreversible pyrrolidino-based first oxidation process [71]. Controlled potential electrolysis (CPE), conducted at a potential more cathodic than the second reduction wave in the first case and more anodic than the first oxidation wave in the second case, affords pure C60 in almost quantitative yields (see Figure 21.5). The generality of the retrocyclopropanation (also termed retro-Bingel reaction) was demonstrated with tris, tetrakis, and pentakis malonate adducts of C60 [70,72]. In these multiple adducts, the number of cyclopropane rings removed could be controlled by the amount of charge transferred. Retrocycloadditions were also successfully carried out with Bingel addends of C70, C76, C78, and C84 [10,70], and the mechanism was investigated with methano C60 derivatives bearing a paramagnetic probe, such as a para nitro phenyl substituted Bingel-like methanofullerene [73]. Based on that study, a dicarbonylic radical derivative was proposed as a reasonable intermediate obtained after cleavage of the cyclopropane ring. Other electrochemically unstable C60 adducts include silyl derivatives [69] and isoxazoles [74]. Finally, a large number of studies have focused on donor–acceptor systems [75], with C60 as the electron acceptor and various groups such as ferrocenes, tetrathiafulvalenes, quinones, porphyrines, phthalocyanines as the donors and both covalent and noncovalent linkages between the two. In this context, electrochemical techniques have served to establish the existence of ground-state interactions between the fullerene and the donor and to determine the HOMO/LUMO gaps, both of which are important factors in materials design for organic electronic applications. Fullerene-based donor–acceptor dyads typically exhibit desirable bandgaps in the range of 1–2 eV. The exohedral reactivity of EMFs appears at first sight to be more complex than that of pristine C60, owing to the diversity and sometimes lower symmetry of the carbon cages [76]. Multiple studies have therefore focused on EMFs with the highly symmetric C80 (Ih) cage, which, provided the

© 2016 by Taylor & Francis Group, LLC

840

Organic Electrochemistry O

O

O

O 1) CPE at –1.5 V vs. Ag wire 4 electrons added 2) Re-oxidation

82%

(a)

O O

O O 1) CPE at –1.5 V vs. Ag wire 6 electrons added

O O

2) Re-oxidation

O O 75% (b)

1) CPE at + 1 V vs. Fc/Fc+ 1.8 electrons removed 2) Re-reduction

(c)

73%

FIgURE 21.5 Retro-Bingel reaction of C60 methano derivatives: (a) Mono C60 Bingel adduct and (b) any C60 Bingel-bis adduct. (From Herranz, M.Á. et al., Eur. J. Org. Chem., 2299, 2004.) (c) Retro 1,3-dipolar cycloaddition of a N-ethyl-pyrrolidino-C60. (From Lukoyanova, O. et al., Angew. Chem. Int. Ed., 45, 7430, 2006.)

endohedral moiety is freely rotating inside the cage, features only two distinct sites for 1,2-double bond addition reactions: the all equivalent [5,6] bonds and the all equivalent [6,6] bonds. In addition, 2 different sites exist for 1,4-addition reactions as well. The chemistry of M@C82 (C2v) EMFs has also been well investigated for a completely different reason. In these compounds, the off-center position of the endohedral atom induces high local strain on specific carbon atoms and a consequent polarization of the cage (see Figure 21.6). The carbon atoms closer to the metal bear a high negative charge density and are thus the target of electrophiles, whereas some carbon atoms on the opposite site of the cage bear positive charges and are preferentially attacked by nucleophiles. Overall, only a limited number of addition isomers have been obtained despite the 24 nonequivalent carbons of these EMFs [8c,76]. Electrochemical studies of such derivatives give insight into how the electronic properties of endohedral fullerenes are modified upon derivatization. 2. Endohedral Metallofullerenes Sharing the C80 (Ih) Cage The icosahedral C80 cage has been derivatized using the typical methods in fullerene chemistry including 1,3 dipolar cycloadditions (leading to 1.2-pyrrolidino adducts) [44,77], Diels–Alder reactions [78],

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Electrochemistry of Fullerenes, Derivatives, and Related Compounds

841

Highest local strain and negative charge density

M

Positive charge density

Positive charge density

FIgURE 21.6 Localization of the local strain and high negative and positive charge density on La@C82 (C2v). (Reprinted from Lu, X. et al., Chem. Commun., 47, 5942, 2011. With permission.)

and Bingel, carbene, or nucleophilic reactions (leading to 1,2-methano adducts) [44–45,79–81]. Often, both [5,6] and [6,6] adducts have been prepared, which offers the opportunity to compare the electronic characteristics of both types of derivatives. In the MNEFs, thermal interconversion of [6,6] adducts into [5,6] ones have been described, which suggests that the [5,6] and [6,6] derivatives are the thermodynamic and kinetic products, respectively [82]. Several of these derivatives have been identified as fulleroids [79–81]. 1,4-Adducts have also been obtained upon silylation [83] and trifluoromethylation reactions [84]. The redox properties of these derivatives suggest that the electron affinity and the electrochemical gap of EMFs can be adjusted through derivatization. Similarly to empty fullerenes, electron-donating addends (adamantylidene Ad, silanes) shift both the oxidation and reduction potentials toward more negative values, whereas electron-withdrawing addends (tetracyanotetrahydrofurane, trifluoromethyl) shift the redox potentials in the reverse direction (see Table 21.8). These displacements are gradual: two Si atoms on La2@C80 and four CF3 groups on Sc3N@C80 shift the reduction potentials approximately twice as much as one Si atom and two CF3 groups, respectively. Nevertheless, in contrast to empty fullerenes, there are no systematic potential shifts for these derivatives due to bond saturation. A significant difference of up to 0.36 V in the oxidation potentials of the [5,6] and [6,6] regioisomers of the pyrrolidino derivatives of M2@C80 has been observed. Remarkably, for the M3N@C80 (M = Er, Lu, Y) derivatives, electrochemical reversibility also depends on the location of the addends: [5,6] adducts normally exhibit reversible reduction processes, whereas [6,6] adducts typically maintain the irreversible reductive behavior of the underivatized MNEFs (see Table 21.8) [44]. As a consequence, electrochemical methods can be useful to determine the location of the addend. Until now, the only exceptions to this rule are [6,6] diphenylmethane [81], tritylpyrrolidino [85] and benzyne adducts [86] of Sc3N@C80, which also exhibit reversible reductive behavior. These results may be another consequence of the significant contribution of the endohedral cluster to the LUMO of Sc3N@C80, which is unique among the MNEFs. The pyrrolidino adducts of MNEFs generally exhibit an addend-based first oxidation process, which is often irreversible. Similarly to C60, CPE at a slightly more positive potential than the first oxidation process results in the removal of the addend [71]. 3. M@C82 (C2v) Endohedral Metallofullerenes The M@C82 (C2v) EMFs are paramagnetic species (see Section II.B.1), and both adducts with open-shell and closed-shell electronic configuration have been described by Akasaka and coworkers [89]. Electrophilic diazo compounds, such as 2-adamantane-2,3-(3H)-diaziridine, and electron-rich dienes, such as pentamethylcyclopentadiene, both lead to paramagnetic doubly bonded monoadducts in a highly regioselective way: the reagent attacks the electron-rich [6,6] bonds closest to the

© 2016 by Taylor & Francis Group, LLC

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TAbLE 21.8 Redox Potentials against Fc/Fc+ of Endohedral Derivatives Sharing the C80 (Ih) Cage Compound

Eox

La2@C80(Ih) Ce2@C80(Ih) La2@C80-Ad [6,6 open] Ce2@C80-Ad [6,6 open] La2@C80-(CClPh) [6,6 open] La2@C80-(CClPh)Ad [6,6 open] (A) La2@C80-(CClPh)Ad [6,6 open] (B) La2@C80-(CClPh)Ad [6,6 open] (C) La2@C80-Dep2SiCH2CHtBp [5,6] (I) La2@C80-Dep2SiCH2CHtBp [5,6] (II) La2@C80-(Mes2Si)2(CH2) [1,4 adduct] La2@C80-(Dep2Si)2(CH2) [1,4 adduct] La2@C80- N-tritylpyrrolidine [6,6] La2@C80- N-tritylpyrrolidine [5,6] Ce2@C80- N-tritylpyrrolidine [6,6] Ce2@C80- N-tritylpyrrolidine [5,6] La2@C80-(CN)4THF [5,6] Sc3N@C80 (Ih) Y3N@C80 (Ih) Er3N@C80 (Ih) Lu3N@C80 (Ih) Gd3N@C80 (Ih) Sc3N@C80-N-tritylpyrrolidine [5,6] Sc3N@C80-N-tritylpyrrolidine [6,6] Sc3N@C80-C(CO2Et)2 [6,6] Sc3N@C80-spirofluorene [6,6] Sc3N@C80-CHPh [6,6 open] Sc3N@C80-(p-C6H13O)2DPM [5,6 open] Sc3N@C80-(p-C6H13O)2DPM [6,6 open] Sc3N@C80-N-ethylpyrrolidine [5,6] Sc3N@C80-C10H12O2 [5,6] Sc3N@C80-C6H4 [5,6] Sc3N@C80-C6H4 [6,6] Sc3N@C80-(Mes2Si)2CH2 [1,4 adduct] Sc3N@C80-(CF3)2 [1,4 adduct] Sc3N@C80-(CF3)4 Sc3N@C80-(CF3)10 Y3N@C80-C(CO2Et)2 [6,6 open] Y3N@C80-N-ethylpyrrolidine [5,6] Y3N@C80-N-ethylpyrrolidine [6,6] Er3N@C80-N-ethylpyrrolidine [5,6] Er3N@C80-N-ethylpyrrolidine [6,6] Er3N@C80-C(CO2Et)2 [6,6 open] Lu3N@C80-N-tritylpyrrolidine [5,6] Lu3N@C80-N-tritylpyrrolidine [6,6]

— — — — — — — — — — — — — 1.01 — 1.02 — — — — — — — — — — —

© 2016 by Taylor & Francis Group, LLC

3

Eox

2

Eox

1

Ered

Ered

Ered

ΔEgapa

Reference

1

2

3

0.95 0.95 0.86 0.89 0.93 1.11 — 0.91 — — — — 0.95 0.63b 0.99 0.62b — 1.09b — — 1.11b — 0.60b 0.60b 1.08b 1.03b 1.08* —

0.56 0.57 0.49 0.47 0.52 0.46 0.45 0.48 0.11 0.13 −0.06 −0.03 0.55b 0.23b 0.56b 0.22b 0.64 0.59b 0.64b 0.63b 0.64b 0.58b 0.35 0.39b 0.56b 0.52b 0.50 —

−0.31 −0.39 −0.36 −0.43 −0.26 −0.48 −0.41 −0.41 −0.50b −0.53b −0.76b −0.70b −0.51b −0.45b −0.55b −0.51b −0.21b −1.26 −1.41 −1.40 −1.42 −1.44 −1.14b −1.06b −1.34 −1.24 −1.48 −1.19

−1.72 −1.71 −1.78 — −1.47 −1.66 −1.63 −1.54 — — — — −1.65b −1.71b −1.75b −1.76b — −1.62 −1.83 −1.83 −1.80 −1.86 −1.54b −1.35b −1.90 −1.96 −2.01 −1.51

−2.13 — −2.33 — −1.67 −1.87 −1.81 −1.89 — — — — −2.19b −2.30b −2.34b −2.25b — −2.37 −2.34 −2.16 −2.26 −2.18 −2.38b −2.35b −2.22 −2.41 −2.40 −1.92

0.87 0.96 0.85 0.90 0.78 0.94 0.86 0.89 0.61 0.66 0.70 0.67 1.06 0.68 1.11 0.73 0.85 1.85 2.05 2.03 2.06 2.02 1.49 1.45 1.90 1.76 1.98 —

[22d] [24b] [79a] [79a] [79b] [79b] [79b] [79b] [83c] [83c] [83b] [83b] [77b] [77b] [77b] [77b] [84] [36a] [44] [44] [45] [38] [85] [85] [45] [45] [80b] [81]

—

+0.52

−1.51*

−1.61*

−2.23*

2.03

[81]

— — — — —

0.62b 0.62b 0.92 0.83b —

0.30 0.25b 0.34 0.42b 0.08b

−1.18b −1.16b −1.11b −1.08b −1.45b

−1.57b −1.54b −1.50b −1.29b —

−2.29b −2.26b −2.21b −2.23b —

1.48 1.41 1.45 1.50 1.53

[44] [44] [86] [86] [83a]

— — — — — — — — — — —

0.60b — — — 0.64b 0.65b 0.64b 0.64b — — 0.64b

0.47b 0.55b 0.86b 0.60b 0.37 0.40 0.34 0.39 0.60b — 0.43b

−1.16b −1.06b −0.84b −1.43 −1.30b −1.40 −1.28b −1.36 −1.43 −1.13b −1.28

−1.65b −1.55b −1.32b −1.88 −1.65b −2.10 −1.63b −2.00 −1.88 −1.42b −2.03

−2.04b −2.03b −2.11b — −2.36b — −2.33b — — −2.43b −2.31

1.59 1.61 1.70 2.03 1.67 1.80 1.62 1.75 2.03

[87b] [87c] [87c] [44] [44] [44] [44] [44] [44] [85] [85] (Continued)

1.71

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Electrochemistry of Fullerenes, Derivatives, and Related Compounds

TAbLE 21.8 (Continued) Redox Potentials against Fc/Fc+ of Endohedral Derivatives Sharing the C80 (Ih) Cage Ered

Ered

Ered

ΔEgapa

Reference

0.27 0.43 0.06

−1.43 −1.52 −1.55

−1.72 −1.73 −2.01

−1.94 −1.99 —

1.70 1.95 1.61

[83d] [83d] [83e]

—

0.08

−1.61

−2.15

−2.58

1.69

[83e]

1.15b 1.10b — —

0.62b 0.59 0.58b 0.59b

−1.45 −1.49 −1.39 −1.40

−1.88 −1.95 −1.83 −1.88

−2.22 −2.32 −2.17 —

2.07 2.08 1.97 1.99

[45] [80a] [88] [88]

Compound

Eox

Lu3N@C80-(SiDep2) [5,6 open] Lu3N@C80-(SiDep2) [6,6] Lu3N@C80-(Mes2Si)2CH2 [1,4 adduct] Lu3N@C80-(Dep2Si)2CH2 [1,4 adduct] Lu3N@C80-C(CO2Et)2 [6,6] Lu3N@C80-CHPh [6,6 open] Gd3N@C80-C(CO2Et)2 [6,6 open] Gd3N@C80-[C(CO2Et)2]2

— — —

— — —

— — — — —

a b

3

Eox

2

Eox

1

1

2

3

ΔEgap = Eox − Ered ; DPM, diphenylmethane; Dep, 2,6-diethylphenyl; Mes, mesityl. Reversibility at 0.1 Vs−1 in the CV experiments. 1

1

metal atom in the former case and the positively charged carbon atoms on the other side of the cage in the latter case as described earlier (Figure 21.5). In contrast to empty fullerenes, these EMFs also react thermally with disilirane reagents, as a consequence of their stronger electron-accepting properties [90]. Several doubly bonded adducts active in electron paramagnetic resonance (EPR) are obtained in this case. In the Bingel–Hirsch reaction with bromomalonate, one EPR-silent singly bonded adduct (La@C82–CBr(COOC2H5)2) is dominant. When malonate is used, a bis-adduct is preferentially formed, La@C82–[CH(COOC2H5)2]2, which retains an open-shell configuration. Under irradiation, direct radical coupling occurs with toluene derivatives, leading to EPR-silent benzyl derivatives. Radical pairing with NO2 has also been observed in benzyne adducts of La@ C82. The paramagnetic adducts retain an electrochemical gap in the range of 0.4–0.6 V, but their reduction and oxidation potentials are shifted with respect to the parent EMF. Studies of this family of compounds (see Table 21.9) confirm that silanes and adamantylidene addends decrease the electron affinity of EMFs. The redox potentials and electrochemical bandgaps of the adamantylidene derivatives are only slightly influenced by the position of the addend [91–93]. As expected, the diamagnetic derivatives exhibit larger electrochemical bandgaps. Their redox potentials are highly dependent on the position of the addend on the carbon cage. A difference in the reduction potentials of up to 0.37 V has been observed for La@C82–CH2C6H5 regioisomers [98]. 4. Endohedral Metallofullerenes Incorporated in Donor–Acceptor Dyads Researchers have started to investigate the highly tunable electrochemical bandgap of EMFs in various donor–acceptor dyads [102]. Like their empty fullerene analogs, EMFs can act as electron acceptors in these assemblies due to their high electron affinity. However, some EMFs also exhibit good electron donor properties (see Section II.B.1) making them potential donors as well [103]. Most interestingly, the role of the donor and acceptor in a Ce2@C80-zinc porphyrin system has been switched by changing only the nature of the solvent [104]. Other donor molecules considered thus far include aza and thiacrown ethers [105], triphenylamine [106a], ferrocene [106b], extended TTF [106b,107], and phthalocyanine [106b]. Covalent linkages via methano bridges or pyrrolidines as well as supramolecular assemblies have been reported [102,108]. These studies have identified electronic interactions between the donor and acceptor groups in the ground state in several of the described systems. In covalently linked triphenyl amine Sc3N@C80 conjugates, for instance [106a], the reduction of the fullerene is shifted toward more positive potentials, while the oxidation of the TPA moiety is shifted toward more negative potentials. The magnitude of the shift depends on the connection between the groups, in this case the substitution on the pyrrolidine. In a very interesting

© 2016 by Taylor & Francis Group, LLC

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Organic Electrochemistry

TAbLE 21.9 Redox Potentials against Fc/Fc+ of M@C82 (C2v) and Some Exohedral Derivatives Compound

Eox

La@C82 (C2v) Gd@C82 (C2v)b Ce@C82 (C2v)b Y@C82 (C2v)b La@C82-adamantylideneb Ce@C82-adamantylidene (I)b Ce@C82-adamantylidene (II)b Gd@C82-adamantylideneb Y@C82-adamantylidene (I)b Y@C82-adamantylidene (II)b Sc@C82-adamantylidene (I)b Sc@C82-adamantylidene (II)b Sc@C82-adamantylidene (III)b Sc@C82-adamantylidene (IV)b La@C82-Me5Cpb La@C82-(Mes2Si)2CH2 (I)b Y@C82-(Mes2Si)2CH2 (I)b Y@C82-(Mes2Si)2CH2 (II)b La@C82-[CH(COOC2H5)2]2b La@C82-CBr(COOC2H5)2 La@C82-CH2C6H5 (2a)c La@C82-CH2C6H5 (2b)c La@C82-CH2C6H5 (2c)c La@C82-CH2C6H5 (2d)c La@C82-CHClC6H3Cl2 (3b)b La@C82-CHClC6H3Cl2 (3d)b La@C82(C6H4)2NO2 b

a b c

2

1.07 1.08 — 1.07 1.01 — — — — — — — — — — — 0.10 — — 0.85 — — — — — — —

Eox

1

0.07 0.09 0.08 0.10 −0.01 0.01 0.02 0.06 −0.02 0.05 0.09 0.09 0.05 −0.04 0.02 −0.07 −0.10 −0.03 0.08 0.38 0.25 0.21 0.17 0.15 0.24 0.25 0.25

Ered

Ered

Ered

ΔEgapa

Reference

−0.42 −0.39 −0.41 −0.37 −0.49 −0.41 −0.42 −0.62 −0.54 −0.43 −0.39 −0.37 −0.42 −0.43 −0.45 −0.50 −0.55 −0.42 −0.32 −0.66 −0.68 −0.95 −0.84 −1.05 −0.91 −0.98 −0.39

−1.37 −1.22 −1.36 −1.34 −1.44 −1.36 −1.35 −1.48 −1.51 −1.37 −1.43 −1.39 −1.40 — −1.71 −1.71 −1.36 — −1.57 −1.31 −1.02 −1.40 −1.42 −1.15 −1.39 −1.07 −1.39

−1.53 — −1.72 −2.22 −1.79 −1.72 −1.74 −1.76 −1.84 −1.70 — — — — −2.22 −1.75 — — — −1.47 −1.21 — −1.74 −1.81 — −1.34 —

0.49 0.48 0.49 0.47 0.48 0.42 0.44 0.68 0.56 0.48 0.48 0.46 0.47 0.47 0.47 0.43 0.45 0.39 0.40 1.04 0.93 1.16 1.01 1.20 1.15 1.23 0.64

[14a] [14a] [14a] [14a] [94] [92] [92] [95] [91] [91] [93] [93] [93] [93] [96] [97] [97] [97] [98] [99] [100] [100] [100] [100] [100] [100] [101]

1

2

3

1 ΔEgap = E1ox/ 21 − Epred . c EPR activity. Compounds 2a–d and 3b–d in [100].

report [109], two covalently linked zinc tetraphenylporphyrin hybrids with La2@C80 and Sc3N@C80 were compared. Studies by nuclear magnetic resonance (NMR) suggested that La2@C80 interacts more with the porphyrin (presumably through strong π–π and electrostatic interactions) than with Sc3N@C80. Electrochemical studies confirmed that the reduction of La2@C80 in the presence of the porphyrin is significantly hindered while that of Sc3N@C80 is almost unaltered. Charge transfer complexes between endohedral fullerenes and organic donor molecules such as N,N,N′,N′-tetramethylp-phenylenediamine (TMPD) have also been reported, in particular (La@C82)• −/(TMPD)• + [110] and (La2@C80)• −/(TMPD)• + [111]. This complexation requires EMFs with not too negative reduction potentials and is temperature and solvent-dependent. The analog complexation with Sc3N@C80 is very weak due to the lower electron affinity of this compound. Upon photoexcitation of the donor, charge-separated states with exceptional lifetimes, often longer than C60 analogs, have been obtained, which means that EMFs could have better electronic properties than empty fullerenes for organic electronic applications. In a Sc3N@C80-ferrocene (Fc) dyad, for instance, after excitation at 388  nm, a charge-separated state (Sc3N@C80)• −/Fc• + with a

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Electrochemistry of Fullerenes, Derivatives, and Related Compounds

lifetime three times longer than an equivalent C60 dyad has been observed [106b,112]. More recently, long-range (45 Å) photoinduced electron transfer and charge-separated states with lifetimes on the order of μs have been achieved for Sc3N@C80-zinc porphyrin dyads [113]. Last but not the least, the first organic bulk-heterojunction solar cell incorporating a clusterfullerene, Lu3N@C80, exhibited a promising open circuit voltage 0.3 V higher than the classical devices using [6,6]-phenyl-C61 butyric acid methyl ester [114]. The short circuit current of this device, however, still needs some improvement.

D.

LARGER CARBON NANOSTRUCTURES: STRUCTURE, DEFECTS, AND IMPURITY EFFECTS

Carbon nanotubes (CNTs) and carbon nano-onions (CNOs) belong to the same family of new carbon allotropes as fullerenes and were discovered almost simultaneously by Iijima [115] and Ugarte [116], respectively, in the early 1990s. CNTs are made of either one (the so-called single-wall carbon nanotubes [SWNTs]) or several graphene layers (multiwall carbon nanotubes [MWNTs]) rolled up into concentric cylinders. They can be metallic or semiconducting depending on their diameter and helicity [117]. CNOs are best described as multishelled fullerenes (see Figure 21.7). In contrast to fullerenes, the solution electrochemistry of pristine CNTs and CNOs is still largely unexplored due to the tedious methods of purification as well as their difficult processibility in solution. Samples of CNTs are highly heterogeneous and contain significant amounts of amorphous carbon, nanographite particles, and metallic impurities [118] that taint the results. In addition, CNTs are essentially insoluble materials, and, consequently, they must either be dispersed using surfactants or functionalized to improve their solubility [119], otherwise studies on CNTs must be performed in the solid state [120]. All of these approaches give valuable information, yet results are obscured by interactions with other tubes (bundling) or other materials in solution (such as surfactants), the substrate (in the solid state), or disruption of the π network (for functionalized nanotubes). Reasonable progress has recently been made by Prato et al., who reported the electrochemical characterization of SWNTs derivatized with pyrrolidines on their sidewalls [121]. The cyclic voltammogram exhibited no discrete peaks, as in the case of fullerenes, but a continuum of diffusioncontrolled reduction current with onset at −0.5 V. This current was related to the electronic density of states of the SWNTs, and from quantum chemical calculations of this density of states, the authors claimed that derivatization alters significantly only the low-lying electronic states but not the overall electronic properties of the nanotubes. A similar behavior was reported for CNTs derivatized with poly(sodium 4-styrenesulfonate) [122] as well as lithium salts of reduced CNTs [119b]. Acid-treated CNTs [123] and CNOs [124], in contrast, exhibit voltammograms with broad anodic and cathodic peaks, attributed to the presence of oxygenated groups on the carbon surface such as carboxylates, ethers, ketones, and lactones. These groups are converted into hydroxyl groups upon reduction.

(a)

(b)

(c) (d)

FIgURE 21.7 Structures of SWNTs, CNOs, and graphene sheets. (a) C60. (b) Ideal single-wall carbon nanotube (SWNT). (c) Ideal carbon nano-onion (CNO). (d) Ideal graphene sheet.

© 2016 by Taylor & Francis Group, LLC

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Organic Electrochemistry

CNTs, however, have received considerable attention as electrode materials [125] due to their unique electrical conductivity, mechanical strength, and high surface-to-volume ratio. They are sometimes considered as the ideal nanometer-sized electrode and have potential applications in electrochemical sensing [126] and energy storage [127]. Numerous studies with various electrochemical probes, including dopamine, ascorbic acid, oxygen, hydrogen peroxide, nicotinamide adenine dinucleotide, norepinephrine, and potassium ferrocyanide [123,128] at CNT-modified electrodes suggested that these materials might enhance voltammetric currents, exchange electrons faster than graphite electrodes, and even exhibit an electrocatalytic effect. The origin of these properties remains highly controversial and is presently strongly debated in the scientific community. Compton et al. have shown that in both MWNT- [129] and SWNT [130]-modified electrodes, electron transfer probably takes place either at the edge-like sites located at the tips of the tubes or at defect sites. In other words, the sidewalls of the tubes might be as much electrochemically inactive as basal planes of graphite, which means that CNT-based electrodes and edge-plane pyrolytic graphite electrodes should in principle exhibit similar properties. The high heterogeneous electron transfer rate and electrocatalytic behavior sometimes observed with CNT-based electrodes actually could be due to some nanographite [131] and metallic impurities [132] contained in the CNT samples or to the oxygenated groups at their tips [133]. Dekker et al., however, have shown that individual SWNTs grown by chemical vapor deposition on a silicon wafer with no tips exposed to the solution could exchange electrons with ferrocenylmethyltrimethylammonium at their sidewalls [134]. Dai et  al. suggested recently that the location of the electron transfer could ultimately depend on the electrochemical probe used for the studies [135]. Given these often contradictory reports, the authors recommend to exercise caution and to take into consideration the structure of the CNTs, their purity, as well as all the chemical treatments they have undergone before claiming any advantage of these materials. These recommendations also apply to graphene and CNOs, which are also being exploited as promising electrode materials [136,137]. Research on CNTs is also directed toward donor–acceptor assemblies for photovoltaic applications. Similarly to endohedral fullerenes, CNTs can act both as electron donors and acceptors, depending on the interaction with photoexcitable molecules [138]. The usual electron donors in the photoexcited state include porphyrins, phthalocyanines, and ferrocene, whereas C60, upon excitation, can accept one electron from SWNTs. As expected, the photoinduced electron transfer rates depend on the diameter of the CNTs as well as their electronic structure. Undoubtedly, electrochemical methods have greatly contributed to the electronic characterization of fullerenes, endohedral fullerenes, and larger carbon nanostructures. The high sensitivity is particularly adept to the small amounts of material that are typically available. The most recent studies have revealed that the electronic properties of EMFs, including their electrochemical bandgaps, can be fine-tuned by changing the encapsulated moiety or the nature of the cage or by adding exohedral groups. This tunability makes EMFs promising candidates for molecular electronics applications. Other carbon nanostructures, including CNTs, are used as electrode materials for promising applications in sensing and energy storage. Where more conventional chromatographic or synthetic methods would fail, electrochemical methods make it possible to isolate and purify fullerenes and related compounds and to prepare specific isomers of multiple derivatives.

III.

PREPARATION, PURIFICATION, AND DERIVATIZATION OF FULLERENES AND RELATED COMPOUNDS bASED ON THEIR REDOX PROPERTIES

A.

ISOLATION OF KINETICALLY UNSTABLE EMPTY AND ENDOHEDRAL FULLERENES

1. Looking for the Missing Empty Carbon Cages A surprising feature of the HPLC traces of the crude mixture of fullerenes obtained from graphite arcing or laser ablation of carbon targets is the absence of empty C72, C74, C80 (Ih), and C88, while larger fullerenes such as C90, C92, and C96 can still be detected. Alford and Diener suggested that the unexpected absence of these empty fullerenes is linked to their very small HOMO–LUMO gap,

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Electrochemistry of Fullerenes, Derivatives, and Related Compounds

as predicted by DFT [139]. Electrochemical studies have confirmed that all the isolated fullerenes feature a HOMO–LUMO gap of 1.0 eV or greater (see ΔEgap in Table 21.3) [12–13], whereas C74, for instance, is predicted to have a bandgap of 0.05 eV and probably a diradical-like electronic structure at ambient temperature [139a]. This carbon cage is thus kinetically unstable and should exist only as an oligomer or polymer, or as a charge transfer complex with a metal like most of the small bandgap fullerenes. Consequently, through electrochemical or chemical reduction, these compounds should in principle get a closed-shell electronic configuration, and therefore increased kinetic stability. Alford and Diener reported the isolation of C74 and fullerenes up to C100 by electrolysis of sublimed raw carbon soot obtained after the arcing process [139a]. At −1.0 V versus Ag/AgNO3 reference electrode both the low and high bandgap fullerenes are reduced to anionic forms and stay in solution due to electrostatic repulsions. Upon reoxidation at 0.4 V, the low bandgap fullerenes return to their polymerized solid state and can be selectively deposited on the working electrode. A second electrochemical reduction eventually yields a solution of low bandgap fullerene anions devoid of C60 and C70. 2. Separating Endohedral Metallofullerenes from Empty Fullerenes Smalley et  al. discovered the first EMF (La@C2n) almost at the same time as C60 after laser vaporization of composite targets made of graphite and lanthanum oxide or chloride [140]. Studying these fascinating compounds has not been an easy task due to their availability in very small quantities, laborious separation from empty fullerenes, and kinetic instability. Fortunately, EMFs are often easier to reduce than the corresponding empty cages (see Section II.B. and Table 21.10), and therefore selective electrochemical or chemical reduction of soot extracts containing a mixture of endohedral and empty fullerenes is a very useful technique to isolate these compounds. Neutral La@C82 and Pr@C82 are radical species. However, their corresponding monoanions are diamagnetic, stable, and very soluble [141]. Along these lines, Akasaka et al. successfully isolated lanthanide-containing metallofullerenes by CPE at 0 V versus standard calomel electrode (about −0.5 V vs. Fc/Fc+) [142]. At that potential (see Table 21.9), the metallofullerenes are selectively reduced and become soluble in a mixture of acetone and CS2, whereas the empty fullerenes remain insoluble. Filtration of the resulting mixture yields a filtrate free of any empty fullerene.

TAbLE 21.10 Comparison of the Redox Potentials of Some Endohedral Metallofullerenes and Empty Fullerenes (in V vs. Fc/Fc+) Compound La@C82 (C2v) La@C82 (Cs) Pr@C82 (C2v) Pr@C82 (Cs) Gd@C82 (C2v) [Li@C60]+ C60 C70 C76 C78 (C2v) C82 C84 a

ΔEgap = E1ox − E1red.

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E1ox (

E1red (

ΔEgapa

Reference

0.07 −0.07 0.07 −0.07 0.09 — 1.26 1.20 0.81 0.95 0.72 0.93

−0.42 −0.48 −0.39 −0.48 −0.39 −0.39 −1.06 −1.02 −0.83 −0.77 −0.69 −0.67

0.49 0.54 0.46 0.55 0.48 — 2.32 2.22 1.64 1.72 1.41 1.60

[14a] [15b] [15b] [15b] [14a] [144] [13] [13] [13] [13] [14a] [13]

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Alternatively, electrochemical or chemical oxidation can also be used to separate or purify some EMFs. Indeed, as a consequence of the electron transfer between the metal and the cage, metallofullerenes are expected to be more easily oxidized than empty fullerenes (see Table 21.10). Solutions of cationic fullerenes can usually be kept over several days in weakly coordinating solvents in the absence of any nucleophile. Bolskar et al. reported the separation of Gd metallofullerenes into three different fractions by selective chemical oxidation [143]. The soot sublimate was first treated with 1,2-dichlorobenzene to remove the soluble empty and metallofullerenes (including Gd@C82). The insoluble Gd@C2n mixture was then treated with the oxidant tris(4-bromophenyl) ammoniumyl hexachloroantimonate (TPBAH) to solubilize and remove the easily oxidizable compounds (Gd@C2n with 72 < 2n < 106). Finally, some of the empty C74 was removed from the insoluble material by oxidation with the stronger oxidant AlCl3. This procedure yielded up to 500 mg of a fraction enriched in Gd@C60, out of 2.5 g of soot sublimate. Other C60 metallofullerenes have turned out to be very difficult to isolate due to their high reactivity toward empty fullerenes, which are also formed in the arc reactor. Only recently, Tobita et al. reported the complete isolation and characterization of Li@C60, as a Li+@C60, SbCl6 − salt [144]. This endohedral fullerene was synthetized by a plasma method in the presence of Li. The carbon deposit was then reacted with TPBAH as oxidant in 1,2-dichlorobenzene at 100°C, yielding the cationic metallofullerene, which proved sufficiently stable for getting a crystal structure. Its ability to form supramolecular complexes with electron donor hosts was also demonstrated [145]. This procedure should be easily extended to other M@C60 metallofullerenes, such as Gd@C60 and La@C60. 3. Looking for the Missing Endohedral Metallofullerenes Since the pioneer work of Smalley et al., it is known that several EMFs are formed during the laser ablation of graphite and lanthanum targets, including La@C60, La@C74, and La@C82 [146]. Until recently, however, only two isomers of La@C82 were easily isolated: the (C2v) and the (Cs) [147] (see Section II.B.1.). This is a consequence of the transfer of three electrons from lanthanum to the cage, which results in paramagnetic and thus potentially reactive EMF species. An important milestone was achieved by Akasaka and coworkers in 2005 with the isolation of La@C74 (D3h) [148a]. The extraction was achieved in 1,2,4-trichlorobenzene, which produces dichlorophenyl radicals upon heating. These radicals react with the paramagnetic fullerenes during the extraction procedure and form diamagnetic adducts with larger electrochemical bandgaps and thus higher stability. In this case, up to six regioisomers of La@C74 –C6H3Cl2 with electrochemical bandgaps in the range of 1.2–1.4 V were obtained [148]. This procedure was also successfully applied for the isolation of  La@C72 (C2) [149], La@C80 (C2v) [150], La@C82 (C3v) [151], and more recently for the elusive M@C60 and M@C70 [152].

B.

SEPARATION OF ISOMERIC CARBON CAGES

1. by Selective Oxidation Sc3N@C80 was first prepared by Dorn and coworkers in 1999 by arcing graphite rods packed with Sc2O3 under a mixture of helium and nitrogen gases [27]. After purification by HPLC of the soluble soot extract, a mixture of two isomers of Sc3N@C80 was obtained: the predominant (Ih) isomer and the less abundant (D5h). The voltammogram of the HPLC-pure sample clearly shows the superposition of waves from both isomers in the anodic region (Figure 21.8a) [36a]. Interestingly, the (D5h) isomer is easier to oxidize than the icosahedral one by about 0.3 V, which allowed Echegoyen et al. to separate the two isomers [36a]. TPBAH that has a redox potential intermediate between those of the two isomers (Figure 21.8b), preferentially oxidizes the (D5h) isomer. It could thus be removed from the mixture by simple filtration over a plug of SiO2. An electrochemical analysis of the pure icosahedral isomer of Sc3N@C80 was then reported as proof (Figure 21.8c). The isolation of the (D5h) isomer was not described, but reductive desorption from the silica could lead to

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Ih+D5h(?) mixture (a) Current/a.u.

Oxidative couple of TPBAH

(b)

Pure Ih (c) 1.2

1.0

0.8

0.6 0.4 0.2 Potential/V vs. Fc/Fc+

0.0

–0.2

FIgURE 21.8 OSWV of (a) mixture of (Ih) and (D5h) isomers of Sc3N@C80; (b) TPBAH; (c) pure (Ih) isomer of Sc3N@C80. (Reprinted from Elliott, B. et al., J. Am. Chem. Soc., 127, 10885, 2005. With permission.)

the recovery of this pure isomer. Milder oxidants such as acetylferrocenium salts [Fe(COCH3C5H4) Cp]+can also be used, which improve the yield of Sc3N@C80 (Ih) recovery [153]. 2. Using the Retrocyclopropanation Reaction The electrochemically induced retrocyclopropanation reaction has found an important application in the isolation of enantiomerically pure chiral fullerene cages. Optically active malonates have been attached to these fullerene cages as chiral auxiliaries. The resulting diastereoisomers were separated by HPLC, and removal of the methano addends by CPE (the retrocyclopropanation reaction) yielded the pure fullerene enantiomers, as confirmed by circular dichroism measurements. This strategy has been successfully applied to C76 [70,73,154] and C84 (D2) [155].

C. PREPARATION OF EXOHEDRAL DERIVATIVES OF FULLERENES 1. by Nucleophilic Substitution Chemical derivatization of fullerenes is usually carried out using cycloaddition, radical addition, metal complex formation, or nucleophilic substitution [156]. The latter involves fullerene anions, which can be easily prepared in aprotic solvents either by chemical or electrochemical reduction. Electrochemical reduction, however, allows for better control of the amount of charge on the carbon cage. In 1993, Kadish et al. reported for the first time the preparation of a mixture of 1,2- and 1,4-dimethyl C60 isomers by the reaction of C602− generated by CPE with methyl iodide in benzonitrile [157]. Different organic halides were later explored, including dibromo substrates and benzyl bromide, leading to substituted methanofullerenes [158] and to a 1,4-(C6H5CH2)2C60 derivative [159a], respectively. Further electrolysis of 1,4-(C6H5CH2)2C60 to generate (C6H5CH2)2C602− by CPE, yielded a mixture of the 1,4;1,4 (Figure 21.9a) and 1,4;1,2 isomers of (C6H5CH2)4C60 [159b]. It should be pointed out that these derivatives are very difficult to prepare by other methods. Interestingly, when carrying out the nucleophilic substitution of benzyl bromide still in benzonitrile but with

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(a)

(b)

FIgURE 21.9 X-ray structures of (a) the 1,4;1,4 isomer of (C6H5CH2)4C60 and (b) fullerooxazole obtained by reaction of electrochemically generated C602− and C603−, respectively, with benzyl bromide in benzonitrile. Hydrogen atoms were removed for clarity. For (b), analysis of the electron density did not allow discrimination between nitrogen and oxygen; however, NMR and mass spectrometry confirmed the presence of an oxazoline heterocycle. These structures were drawn using Mercury and crystallographic information files QAWRUB [159b] and PBCA [160a].

C603− instead of C602−, Gao et al. recently obtained a new and unexpected C60 cycloadduct: a fullerooxazole (Figure 21.9b) [160a]. Surprisingly, it turned out that the solvent was the source of nitrogen, whereas the oxygen atom came from traces of water in the reaction mixture. This reaction was reproduced in the absence of benzyl bromide [160a], with various substituted benzonitriles [160b] and a methano derivative of C60 (C61HPh) to reveal the regioselectivity [160b]. 2. Using the Retrocyclopropanation Reaction Since malonate methano addends can be easily and selectively removed in the presence of pyrrolidino addends, they have also been used as blocking groups on the carbon sphere to get better control of the regiochemistry of multiple additions. Using the orthogonal transposition method introduced by Kräutler et al. in 1997 [161], four di(ethoxycarbonyl)-methano groups can be added to the equatorial position of C60. Due to steric reasons, further addition of pyrrolidine groups will occur at the two poles of the carbon sphere. Subsequently, some of the methano groups can be selectively removed by retrocyclopropanation. Along these lines, the selective formation of an improbable (e,e,e,e)-tetrakis-[60] fullerene derivative was recently reported. The compound bears two pyrrolidine groups trans-1 to each other and two cyclopropane rings also trans-1 to each other (Figure 21.10) [162]. The reductive behavior of this compound was surprisingly reversible (see Figure 21.11), which suggests that the remaining two cyclopropane rings cannot be easily removed. 3. Fine-Tuning the Reactivity of Endohedral Fullerenes So far, the chemistry of endohedral fullerenes has not been studied as extensively as that of the empty fullerenes, mainly due to the small amounts of most of the compounds that can be obtained from the soot. However, some very interesting studies have suggested that the chemical properties of these molecules can be fine-tuned by redox treatment prior to the reaction. Akasaka et  al. compared the reactivity of M@C82, (M@C82)+, and (M@C82)− (M = Y, La, Ce) toward 1,1,2,2-tetramesityl-1,2-disilirane [163]. This nucleophilic reagent has been frequently used as a probe of fullerene reactivity. The cationic and anionic fullerenes were generated by oxidation with

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Electrochemistry of Fullerenes, Derivatives, and Related Compounds N

N 1) anthracene, 180° C 2) diethylbromomalonate/ diazabicycloundecene, CH2Cl2, room temperature 3) 195° C R R Orthogonal transposition

R R R R

R = CO2Et

R

pyridine glycine, paraformaldehyde, o-dichlorobenzene, reflux

R

N

N

R R R R

R R

Two successive 1,3-dipolar cycloaddition

R R

1) CPE at – 1.9V vs. Ag wire 6 electrons added 2) Reoxidation R

R R R

Retro-Bingel reaction N

N

N 3%

N 18%

FIgURE 21.10 Synthesis of an (e,e,e,e)-tetrakis-[60] fullerene derivative using the orthogonal transposition strategy and a selective retrocyclopropanation. (Reprinted from Ortiz, A.L. and Echegoyen, L., J. Mater. Chem., 21, 1362, 2011. With permission.) 1.4 1.2 1.0

Current (μA)

0.8 0.6 0.4 0.2 0.0 –0.2 –0.4 –0.6 –1.0 Potential/V vs.

–1.5

–2.0

Fc/Fc+1

FIgURE 21.11 CV of the (e,e,e,e)-tetrakis-[60] fullerene derivative obtained in CH2Cl2 + 0.1 M (n-Bu)4NPF6 (scan rate 0.1 Vs−1).

TPBAH and reduction with sodium thiomethoxide, respectively. As summarized in Table 21.11, the cationic fullerenes exhibited enhanced thermal reactivity, when compared to their corresponding neutral species. In contrast, the anionic fullerenes did not react at all. These observations were correlated to the redox potentials of the different species. Easily reducible fullerenes are electrophilic and react easily with disilirane, whereas less reducible fullerenes do not react (see Table 21.11). Conversely, more nucleophilic endohedral fullerene species should react more easily with electrophiles. A very recent study by Echegoyen et al., involving trimetallic nitride cluster fullerene dianions and benzyl bromide as electrophile confirmed this fact [80a]. In that report, the dianions of Sc3N@C80 (I h) and Lu3N@C80 (I h) were generated by CPE at −1.6 V versus silver wire reference electrode. DFT calculations, as well as the electrochemical properties of these two compounds (see Table 21.10), suggested that Lu3N@C80 dianion is more nucleophilic than its scandium analog. Indeed, only the dianion of Lu3N@C80 reacted with benzyl bromide, leading

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TAbLE 21.11 Redox Potentials against Ferrocene/Ferrocenium Fc/Fc+ of Some Endohedral Metallofullerenes and Reactivity toward Disilirane and benzyl bromide Compound

E1ox (V )

E1red (V )

Reaction with Disilirane

La@C82 (C2v) Y@C82 (C2v) Ce@C82 (C2v) (La@C82)+ (C2v) (Y@C82)+ (C2v) (Ce@C82)+ (C2v) (La@C82)− (C2v) (Y@C82)− (C2v) (Ce@C82)− (C2v) (Sc3N@C80)2− (Ih) (Lu3N@C80)2− (Ih)

0.07 0.10 0.08 1.07 1.07 1.08 −0.42 −0.37 −0.41 −1.56a −1.80b

−0.42 −0.37 −0.41 0.07 0.10 0.08 −1.37 −1.34 −1.41 −0.83a −2.26b

Yes (80°C) Yes (80°C) Yes (80°C) Yes (r.t.) Yes (r.t.) Yes (r.t.) No No No

a b

Reaction with benzyl bromide

Reference

No Yes (r.t.)

[163] [163] [163] [163] [163] [163] [163] [163] [163] [164] [164]

2nd and 3rd reduction waves of Sc3N@C80. 2nd and 3rd reduction waves of Sc3N@C80.

to the open methano adduct (fulleroid) at the junction between two six-member rings ([6,6] ring junction). Interestingly, nucleophilic addition to Sc3N@C80 (I h) becomes possible if the trianion is prepared [80b].

D.

SOLUTION PROCESSING OF LARGER CARBON NANOSTRUCTURES

Interestingly, redox-based methods have not only advanced the field of nanocarbon as it applies to the fullerene families, but it is also playing a crucial role in advancing the chemistry of nanotubes and graphene. For example, reduction with alkali metals breaks the bundles of CNTs in polar solvents without the need of surfactants, thanks to the electrostatic repulsion between the charges introduced on the tubes [119b,c], and allows their further functionalization [164]. This process has allowed investigators to report the electrochemical characterization of individual semiconducting SWNTs in solution. Electrochemical exfoliation of graphite is also becoming one of the cleanest and easiest methods to obtain dispersions of graphene layers. In initial reports, a high potential was applied between two high purity graphite electrodes immersed in an ionic liquid or an ionomer solution [165]. Exfoliation of graphite occurred at the anode and formed a black precipitate that could be redispersed in polar aprotic solvents. Very recently, exfoliation was also achieved by applying a high negative voltage to a carbon assembly in propylene carbonate, followed by addition of LiCl and sonication [166]. The so-called electrochemical unzipping of CNTs proposed by Pillai et al. is an alternative method to produce graphene nanoribbons [167]. An anodic potential is applied first to break the sp2 carbon bonds and to open the tubes. Then, application of a cathodic potential yields graphene materials with smooth edges.

IV. CONCLUSION Information about the electronic properties of fullerenes obtained from their electrochemical characterization can be successfully exploited to prepare, isolate and/or purify fullerenes that would be otherwise very challenging to obtain. In particular, a series of highly reactive empty fullerenes as well as endohedral fullerenes have been extracted from soot mixtures by either electrolysis or redox processes. The main advantage of these methods is that they produce large amounts of these species

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in short periods of time, as opposed to the labor-intensive, time-consuming separation by HPLC. Also, fullerene anions can be easily produced by CPE and reacted with various substrates, leading to new adducts and cycloadducts. These derivatives are usually not accessible by commonly used fullerene derivatization procedures. Electrochemical or chemical reduction can also help to process CNT and graphene layers in solution. Electrochemically induced retrocycloaddition reactions have found some important applications in the isolation of enantiomerically pure chiral carbon cages, as well as in the preparation of derivatives with complicated addition patterns.

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Aliphatic and Aromatic Hydrocarbons Reduction Jürgen Heinze

CONTENTS I. II.

Introduction .......................................................................................................................... 861 Formation of Radical Anions, Dianions, and Polyanions .................................................... 862 A. Experimental Aspects ................................................................................................... 862 B. Redox Properties ........................................................................................................... 863 C. Electron Transfer Kinetics ............................................................................................ 873 III. Chemical Reactions of Electrogenerated Anions ................................................................. 874 A. Homogeneous Electron Transfer................................................................................... 874 B. Electrophilic and Related Reactions ............................................................................. 876 1. Protonation ............................................................................................................. 876 2. Alkylation .............................................................................................................. 878 3. Acylation ................................................................................................................ 879 4. Addition of CO2 ..................................................................................................... 879 C. Reductive Coupling....................................................................................................... 880 D. Intramolecular Reactions .............................................................................................. 881 1. Conformational Changes ....................................................................................... 881 2. Bond-Breaking and Bond-Making Reactions........................................................ 881 References ......................................................................................................................................884

I. INTRODUCTION Hydrocarbons are the simplest organic compounds consisting only of carbon and hydrogen atoms. There are two basic structures, saturated (aliphatic) and unsaturated (olefinic or aromatic) systems. Saturated compounds solely contain σ-bonds within the carbon skeleton, whereas unsaturated compounds at least contain one π-bond. The electrochemical reduction of such species is exclusively restricted to olefinic or aromatic compounds. The reason is that aliphatic hydrocarbons have extremely low electron affinities that render their reduction impossible, despite a gain of solvation energy within the stability limits of conventional solvent–electrolyte systems. Thus, electrochemical data involving both thermodynamic and kinetic parameters of hydrocarbons are available for only olefinic and aromatic π-systems. The reduction of aromatics in particular had already attracted much interest in the late 1950s and early 1960s. The correlation between the reduction potentials and molecular orbital (MO) energies of a series of aromatic hydrocarbons was one of the first successful applications of the Hückel molecular orbital (HMO) theory, and allowed to develop a coherent picture of cathodic reduction [1]. The early research on this subject has been reviewed several times [2–4].

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Later, during the golden age of mechanistic chemistry, interest focused on the elucidation of reaction paths of cathodically generated species, including disproportionations [5] and chemical follow-up processes [6]. Techniques were developed and improved to measure both the kinetics of coupled chemical reactions [7] and thermodynamic parameters such as redox potentials of higher charged aromatics [8]. The discovery of the redox properties of conducting polymers in 1978 and later of the fullerenes [9] gave new impetus to the electrochemistry of unsaturated hydrocarbons, initiating extended studies on the redox properties of these species (Chapter 21) [10]. However, the new challenges also document that the period of basic research on these systems has been concluded and now trends to applied sciences dominate.

II. FORMATION OF RADICAL ANIONS, DIANIONS, AND POLyANIONS A.

EXPERIMENTAL ASPECTS

Thermodynamic reduction potentials of numerous aromatics were first measured by Hoijtink and van Schooten in 96% aqueous dioxane, using polarography [11]. These fundamental works were decisive tests of the HMO theory, showing that the polarographic half-wave potentials vary linearly with the HMO energies of the lowest unoccupied molecular orbitals (LUMOs) of the hydrocarbons [1]. Hoijtink et al. had already noticed that most aromatics can be further reduced to their respective dianions [12]. They proposed a two-step reduction scheme, in which both redox potentials E10 and E20 are on average separated by about 500 mV: E10

−•  ⇀ A↽ A E20

2−  ⇀ A− • ↽ A

(22.1) (22.2)

However, their careful analysis also showed that most of the dianions were not stable in the polarographic or voltammetric time scale, and even less so after bulk electrolysis, and underwent follow-up reactions with water or other electrophilic impurities, details of which are discussed in Section III. Aprotic solvents such as acetonitrile [13,14] or dimethylformamide [15–17] considerably improved the stability of the radical anions but normally had little effect on the reactions of the more basic dianions [17,18]. The increased irreversibility of the dianion formation is probably due to the ability of dianions to abstract protons even from the solvent, or, by Hofmann elimination, from the tetraalkylammonium salts that are common supporting electrolytes in aprotic solvents [19]. Progress in electrochemical instrumentation soon stimulated the application of more elaborated measurement methods than simple dc polarography, which facilitated studies of heterogeneous kinetics and detection of follow-up reactions of the electrogenerated species. Thus, conclusions originally drawn from the shape and height of polarographic curves have been amply confirmed by straightforward diagnostic criteria in cyclic voltammetry [7,20,21], nowadays the standard method for mechanistic studies in organic electrochemistry [22]. A fundamental improvement in the facilities for studying electrode processes of reactive intermediates was the purification technique of Parker and Hammerich [8]. They used neutral, highly activated alumina suspended in the solvent electrolyte system as a scavenger of spurious impurities. Thus, it was possible to generate a large number of dianions of aromatic hydrocarbons in common electrolytic solvents containing tetraalkylammonium ions. It was the first time that such dianions were stable in the time scale of slow-sweep voltammetry. As the presence of alumina in the solvent– electrolyte systems may produce adsorption effects at the electrode, or in some cases chemisorption and decomposition of electroactive species, Kiesele constructed a new electrochemical cell with an

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integrated alumina column for electrochemical studies under superdry conditions [23]. Such sophisticated methods made it possible to generate reversibly polyanions up to octaanions of aromatic and olefinic hydrocarbons [24]. Further progress in stabilizing highly charged anions was achieved by the application of unconventional solvents such as ammonia or dimethylamine (DMA) at low temperatures. Using these solvents, it was possible to observe the reversible generation of supercharged anions at low scan rates in voltammetric experiments [25,26].

B. REDOX PROPERTIES For purpose of description, the electrochemistry of unsaturated hydrocarbons may usefully be classified in three categories: benzenoid, nonbenzenoid, and olefinic hydrocarbons, each of which exhibit characteristic properties upon reduction. Benzenoid hydrocarbons have been studied in greatest detail [2,4,27]. In aprotic solvents, they can be reversibly reduced to their respective anions without difficulties. Even the electrochemical reduction of benzene in dimethoxyethane has been described [28]. In many cases, the electrogenerated radical anions are stable enough to allow the simultaneous measurements of their ESR spectra [29]. Generally, it can be stated that the positions of the thermodynamic reduction potentials depend on the magnitude of the π-systems. They are shifted to more positive potentials with the increase of the conjugation lengths or the rise in numbers of π-electrons (Table 22.1). The main reason for these changes is the increase of the electron affinity as a function of the π-electron structure [30]. A very striking feature of benzenoid hydrocarbons is the excellent correlation between their thermodynamic reduction potentials and the predictions of semiempirical π-electron theories, especially of the Hückel approximation (HMO). As the thermodynamic reduction potentials are a measure of the electron affinities of the respective compounds, they can be compared with the theoretically calculated energies in the simple MO picture, in which additional electrons have been added to antibonding MOs of the π-systems. Therefore, assuming that the solvation energies for a series of aromatic hydrocarbons are constant, there should be a linear correlation between the thermodynamic reduction potentials (half-wave potentials E1/2) and the calculated energies mm +1 of the lowest unoccupied MO (LUMO) in units of an effective β in the HMO approximation (Equation 22.3): E1/ 2 = −bmm +1 + C

(22.3)

where b corresponds to an effective value of the resonance integral β C is a constant within a series of hydrocarbons Independent voltammetric and polarographic measurements carried out in different solvents such as 2-methoxyethanol, 96% dioxane and dimethylformamide (DMF) confirm the relationship through excellent linear correlations with slopes b of approximately 2.40 [1,15,31]. Later on, Fry [32] applied a modified HMO approximation (ω-technique) and improved the validity of Equation 22.3 by introducing nonbenzenoid polycyclic alternant and nonalternant hydrocarbons, annulenes, cyclophanes, and polyenes. Recently, based on the density functional theory and well-developed computational solvation methods, Fry and Davis succeeded in computing absolute reduction potentials that then were linearly correlated with experimental data (R2 = 0.9981) [33]. Rather surprisingly, the differences in half-wave potentials of hydrocarbons from one solvent to another are very small. This constancy in energy values as well as slopes of correlation lines in widely varying solvents and supporting electrolytes implies that solvation energies, provided that they are not small, change in the same way from system to system. As already observed by Hoijtink [12], nearly all benzenoid hydrocarbons can be reduced to their respective dianions—only benzene and naphthalene are exceptions. In all experiments, these

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TAbLE 22.1 Redox Potentials E 0 for the Reduction of Aromatic and Olefinic Hydrocarbons Compounds Benzene Naphthalenea Anthracenea Tetracenea Pentacenef Chryseneb Coroneneb 1,2-Benzanthraceneb Phenanthrenea Triphenylenea Pyrenea Biphenylc p-Terphenyla o-Terphenyl Quaterphenyla Stilbenea Bianthryla Biphenanthryla Decacyclenec 2,2′-Distyryl-biphenyla 4,4′-Distyryl-biphenylc p-Oligophenylenevinylene (n = 1)a (n = 2)a (n = 3)a c Acepleiadylene COTd Azulenea 1,3,5-Tri-tert-butylpentalened Heptalenef [12]Annulenee 1,7-Methano[12]annulenef 15,16-Dihydropyrene([14]annulene)g [16]Annulenef [18]Annulenef Heptalenef C60h a

a

b c d e f g h

0 ER/R − (V)

−3.42 −2.53 −2.04 −1.55 −1.28 −2.27 −1.99 −1.92 −2.49 −2.42 −2.29 −2.68 −2.40 −2.62 −2.28 −2.26 −1.92 −2.35 −1.74 −2.12 −2.30 −2.00 −1.86 −1.85 −1.85 −1.71 −1.62 −1.41 −1.41 −1.35 −1.42 −2.22 −1.19 −1.52 −1.41 −0.98

ER0− /R2− (V)

ER02− /R3− (V)

ER03− /R4− (V)

References

— — −2.64 −2.15 −1.87 −2.77 −2.63 −2.58 −3.13 −2.97 −2.91 −3.18 −2.70 −2.72 −2.455 −2.72 −2.14 −2.51 −2.14 −2.21 −2.52 −2.24 −1.97 −1.91 −2.51 −1.84 −2.6 — −2.11 −2.00 −1.74 −3.01 −1.48 −1.86 −2.11 −1.37

— — — —

— — — —

— — — — — — — — — — −2.82 −3.23 −2.35 −2.81 −3.06 − −2.79 −2.46 −3.11 —

— — — — — — — — — —

[26,28] [26] [26] [26] [36] [8b] [8b] [8b] [26] [26] [26] [10a] [26] [103a] [26] [26] [26] [26] [22] [205b] [205b] [24a] [24a] [24a] [37] [48] [26] [70] [69] [60] [69] [54] [57] [59] [69] [10c]

— — — — — — — — −1.87

−3.17 — −2.68 (−2.88) −3.13 — − −3.10 −2.88 −3.14 — — — — — — — — — −2.35 (−2.85, −3.26)

Cyclic voltammetry was performed at a Pt electrode with solutions of 10−3 to 10−4 M in substrate. All potentials are expressed in V versus Ag/AgCl; solvent, dimethylamine-TBABr; temperature between −40°C and −65°C. DMF-Me4NBr, potentials versus Ag/AgCl (corrected from SCE). THF-NaBPh4 or LiBPh4, potentials versus Ag/AgCl, in parentheses potentials for penta- and hexaanion formation. COT = cyclooctatetraene; ACN-TEAP, potentials versus Ag/AgCl (corrected from SCE). THF-TBAClO4, potentials versus Hg pool. DMF-TBAClO4, potentials versus Ag/AgCl (corrected from SCE). DMF-TBAClO4, potentials versus Ag/0.1 M AgNO3. Toluene/ACN-TBAPF6, potentials versus Fc/Fc+ at –10°C, in parentheses potentials for penta- and hexaanion formation.

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second reduction steps appear approximately −0.55 ± 0.10 V more negative than the first reduction, provided that the supporting electrolytes used were tetraalkylammonium salts. Therefore, these reduction potentials were also correlated with the LUMO energies of the HMO model [3]. It was suggested that the energy difference of 0.55 eV corresponds to the repulsion energy between both electrons in the LUMOs of the dianions [1], despite the differences in their structures. On the other hand, quantum-mechanical calculations show that the repulsion energies are much larger. Dewar has calculated values in the range of 5 eV [34]. The discrepancy between experiment and theory results from the fact that ion-pairing effects and solvation influences have been neglected in the calculations. Experimental data clearly reveal that counterion effects that efficiently shield the negative excess charge have in the past been underestimated and are considerably stronger for di- and polyanions than for radical anions [35–37]. In polar solvents such as DMF or acetonitrile, the interaction increases in the order But4N < Prop4N < Et4N < K < Na < Li and, consequently, reduction potentials shift in a positive direction [6,38]. Obviously, ion pairing is greatest for the small lithium ions, in agreement with the prediction of Born’s equation [39]. Surprisingly, in the case of Me4N+ as counterion, the ion-pairing effect is significantly diminished [40]. In the case of solvents with low dielectricity constants, the pattern is different, and ion pairing becomes dominant as the radius of alkali cations increases [37,41,42]. The reasons for this behavior have not yet been studied in detail, but it has been proposed that in ethereal solvents the solvation of small cations remains stronger than that of larger ones, and therefore ion pairing of potassium should be more pronounced than that of lithium. Ion-pairing effects may considerably influence disproportionation mechanisms that involve homogeneous redox reactions of anions to their respective dianionic and neutral species (Equation 22.4) [43]: 2A− • ⇌ A 2− + A

(22.4)

Disproportionation mechanisms have been proposed for protonation reactions and intramolecular rearrangements (see Sections III.B and III.C) [44]. Their prominent feature is that follow-up processes at the level of the dianion can already take place at potentials corresponding to radical anion formation. In order to evaluate data for disproportionation reactions, it is necessary to know the value of the disproportionation equilibrium constant: KD =

[ A 2− ][ A] [ A− • ]2

RT ln K D = E20 − E10 nF

(22.5a)

(22.5b)

This can be determined from the difference in reversible potentials of the couples A/A− • and A− •/A2− (Equation 22.5). In the case of benzenoid aromatics, KD values range between 10 −9 and 10 −13, provided that tetraalkylammonium salts have been used as supporting electrolytes [8b]. In solvents of low dielectricity constant, additional effects are observed, showing influences of the supporting electrolyte concentration and of the nature of the cations [6]. In the tetraalkylammonium series, the strongest (contact) ion pairs are formed by Et4N+, and KD is largest for that cation [8b,45]. Drastic changes in the disproportionation constants occur when alkali cations are used instead of tetraalkylammonium ions. Typically, the potentials of the radical anion formation are less affected than that of the dianion formation. In the presence of alkali cations, ΔE0 shifts may reach values of more than 600 mV, which correspond to an increase in the K constant of more than 10 orders of magnitude [43]. However, in a recent publication Fry claimed that disproportionation is driven by solvation, not ion pairing [43c]. It might be that the strong solvation effects at the dianion level have been overlooked.

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TAbLE 22.2 Structure Formula of Annulenes and Cyclic Conjugated Systems (1–17)

2

1

5

3

6

10

4

7

11

8

9

12

13

+ –

+ 14

15

16

– 17

The electrochemistry of nonbenzenoid hydrocarbons (see Table 22.2) has attracted much interest because their structures offer unusual insights into π-electron systems that undergo geometric changes upon reduction and obey both the Hückel 4n and 4n + 2 rules. The most widely studied examples are cyclooctatetraene (COT, 1) and its derivatives. In such conventional aprotic solvents as DMF, dimethyl sulfoxide (DMSO), or acetonitrile containing tetraalkylammonium salts, two distinct one-electron reduction waves are observed at approximately −1.64 and −1.80 V versus SCE, with ΔE 0 separations varying from −130 to −240 mV [36,46–49]. In THF and NH3, this separation reduces further [50–52], and in the presence of alkali salts [52,53] even two-electron reduction waves with positive ΔE0 differences were obtained, indicating large disproportionation constants. The unusually small separation of the two redox steps in comparison to the data of benzenoid aromatics was ascribed to the fact that the planar COT dianion forms a 4n + 2 π-electron system that is stabilized by its gain in Hückel resonance energy (Figure 22.1) [54]. Obviously, strong ion-pairing effects additionally favor the formation of the dianion [52,53]. On the other hand, it was argued that a negative shift of the first reduction step, leading to small ΔE 0 values, is caused by energy requirements accompanying the transition from the tube-shaped to the planar molecule. This is supported by electrochemical results obtained with methyl-substituted COT derivatives. Thus, all methyl-substituted derivatives due to a steric barrier are harder to reduce than 1 itself, and in 1,2,3,4-tetramethylcyclooctatetraene (2) reduction occurs at the extremely negative potential of −3.6 V (HMPA vs. SCE) [55]. In the case of phenyl-substituted COTs, the steric demand is by and large energetically compensated through the planarization of the system and the subsequent resonance interaction between the phenyl moieties and the COT ring. Therefore, the reduction potentials are similar to that of unsubstituted 1. In addition, the resonance-stabilizing effect upon the dianion is sufficient to shift the second redox potential E20 positively to E10 , thus producing a single

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1 μA

0.5 μA

–1.0

–2.0 –1.5 E (vs. Ag/AgCl)(V)

FIgURE 22.1 Experimental cyclic voltammograms of the reduction of COT in THF/0.1 M TBAPF6, ν = 100 mV/s, ΔE0 = 170 mV; obtained from simulations: (at the top) activated (after purification of surface by polishing) Pt electrode, k1 = 1.3 × 10 −3 cm/s, k2 = 10 −2 cm/s; (below) passivated Pt electrode k1 = 3.0 × 10 −4 cm/s, k2 = 10 −2 cm/s. The spike results from homogeneous comproportionation reactions between the dianion and the neutral species (see Section III.A).

two-electron reduction wave [48]. It is interesting to note that the electrochemistry of the doubly decked [8]annulene (= [22](1,5)-cyclooctatetraenophane) also indicates some sterical strain [56]. The formation of the monoanion occurs at a potential of −2.43 V (DMA vs. Ag/AgCl). It can be reduced up to its tetraanion at a potential of −3.22 V. The redox behavior of a number of higher annulenes than 1 has been studied during the 1970s and 1980s. The evaluation of their electronic properties has attracted much interest because it offers a good comparison of the fundamental differences between 4n and 4n + 2 π-systems. As predicted by the simple MO theories, the transfer of an electron to a [4n]annulene should be energetically favorable because the electron is inserted into a low-lying nonbonding orbital. The further reduction to the dianion leads to a stabilized (aromatic) 4n + 2 π-system, which therefore should also be easily accessible. In contrast, during the reduction of an aromatic [4n + 2] annulene electrons are injected into a LUMO with high energy, and in the second reduction step an unstable 4n dianion is generated. Consequently, the reduction of [4n]annulenes should be observed at relatively positive potentials with small ΔE0 separations for the dianion formation, while the reduction of the [4n+2]annulenes should occur at more negative potentials with large ΔE0 separations for the dianion formation. This is exactly what is observed (see Table 22.1). Although benzene [28], the classical Hückel aromatic with 4 × 1 + 2 = 6 π−electrons, is reduced at −3.42 V (vs. Ag/AgCl), the reduction of the [8]annulene COT occurs at −1.71 V. Similarly, the [16]annulene (3) [57,58] is more easily reduced than the corresponding [18]annulene (4) [59], although the reduction of the larger π-systems should be more favorable for electrostatic reasons. The only exception is [12]annulene where the ΔE0 separation is quite large [60]. For larger annulenes, even the reduction to tetraanions is possible [37,61,62]. This is especially favorable for neutral 4n + 2 species, which, after the injection of four electrons, reform a stabilized 4n + 2 π-system. On the other hand, the existence of four excess charges in one molecule gives rise to strong electron repulsion. It can be shielded only by ion pairing that drives the formation of the tetraanion to a potential sufficiently positive for reduction to occur within the stability range of the

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solvent–electrolyte system. Thus, the electrochemical reduction of acepleiadylene (5) to its tetraanion was only achieved in THF in the presence of LiBPh4 as supporting electrolyte [37]. Attempts to generate the same species in the presence of tetrabutylammonium ions were unsuccessful. Several studies have been published concerning the problem of homoaromaticity and its influence on reduction potentials. The term homoaromaticity is applied to cyclic resonance-stabilized systems in which a carbon–carbon double bond is replaced by a cyclopropane ring, or a saturated carbon atom is introduced into the conjugated chain. Paquette et al. investigated the reduction of cis-bicyclo-[6.1.0]-nona-2,4,6-triene (6) and similar derivatives [63]. They found that it is reduced (−2.59, −2.79 V) more negatively than 1 but more easily than cyclooctatriene (7) (−2.77 V), thus proving a degree of homoconjugative stabilization in the radical anion of 6. Analogously, dibenzonorcaradiene (8) is easier to reduce than biphenyl, but harder to reduce than phenanthrene [64]. The electrochemical reduction of 1,6-dimethylbicyclo[4.4.1]undeca-2,4,7,9-tetraene (9) was also interpreted in terms of homoconjugation in the anions. The reduction appears to proceed via an ECE scheme in which the initial reduction produces an unstable radical anion that undergoes a structural change, producing another radical anion with a conjugatively stabilized π-system. The second electron transfer occurs more easily than the first one, thus producing a typical two-electron reduction [65]. Of the condensed nonbenzenoid aromatics, azulene (10) has gained greatest popularity [66]. Although it is an isomer of naphthalene, its electrochemical behavior differs markedly from its benzenoid counterpart. Azulene is reversibly reduced at −1.62 V (vs. Ag/AgCl), while the reduction of naphthalene takes place at −2.53 V [26,67]. Furthermore, the anion of azulene is extremely stable against the attack of protons [67]. The bicyclus octalene (11) has 14 π-electrons. 1H-NMR data show that it possesses a nonplanar structure with polyolefinic properties and therefore resembles its monocyclic relative COT. Cyclic voltammetry reveals that it is reversibly reduced to the radical anion at E1/2 = −1.67 V and, rather surprisingly, in a three-electron process at E1/2 = −1.70 V to its stable tetraanion [68]. This unusual behavior can be only explained by assuming that the energy gain from delocalization in the planar 18 π-system is higher than torsional and electronic repulsions. Other bicyclic systems such as heptalene (12) [69] and pentalene (13) [70] can also be reduced to their respective anions. Nevertheless, in comparison with equivalent monocyclic π-systems such as 1, their reduction behavior is still not properly understood. A further group of nonbenzenoid aromatics is the series of odd-membered cations and anions such as cyclopropenium (14) and tropylium cations (15) as well as cyclopentadienyl (16) and cyclononatetracenyl anions (17). Regarding the arguments for the properties of Hückel-like 4n + 2 π-systems, all these molecules should be energetically stabilized. Obviously, this is not fulfilled in all cases. The tropylium cation (15) can be reduced in a one-electron step to the tropyl radical even at E = +0.06 V versus SCE [71]. The radical is unstable and rapidly dimerizes to bitropyl. The heptaphenyl tropylium radical is stable on the voltammetric time scale, but decays slowly producing a dimeric species involving coupling via the phenyl groups [72]. Similarly, 2,3-diphenylethyl-cyclopropenyl is irreversibly reduced in CH3CN at Ep = −0.04 V. On the other hand, the reduction of trimethylcyclopropenyl occurs at −1.32 V, which is in better agreement with the prediction of the Hückel theory [73–75]. Otherwise, relatively positive reduction potentials seem to be typical for most carbocations [76]. The easy reducibility is probably caused by the excess positive charge. The discovery of the metal-like properties of conducting polymers has once again focused attention on the oxidation and reduction characteristics of aromatic systems. It turns out that most of these conducting materials consist of chainlike connected carbocyclic or heterocyclic aromatics [77–79]. The simplest molecules in these series are dimers, followed by oligomers of increasing chain length and polymers. As the current–voltage curves of polymers are difficult to interpret, quantitative information on the redox properties of such systems was preferentially obtained from reduction experiments with dimers and defined oligomers. Redox data on dimers are available for biphenyl [10a,26], bianthryl [43b,80], biazulenyl [81], bicyclooctatetraenyl [82], and dimeric [14]annulenes [83]. All species can be reduced to at least their respective dianions. In the case of bianthryl and bicyclooctatetraene, even the formation of tetraanions has been observed. In general, the observed

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203 K 2 V/s

5 μA

233 K 100 mV/s

0.5 μA

268 K 100 mV/s 0.2 μA

–3.5

–3.0

–2.5

–2.5

–1.5

E [vs. Ag/AgCl](V)

FIgURE 22.2 Cyclic voltammograms for the reduction of biphenyl, terphenyl, and quaterphenyl in dimethylamine/0.1 M TBABr. Redox potentials have been published in Reference 26.

redox potentials of the monoanion formation differ significantly from those of the monomeric parent compounds and shift to less negative values. This is evidence of conjugative stabilization in the charged oligomeric unit [84]. On the other hand, large ΔE 0 separations between the redox potentials of the respective mono- and dianions indicate strong electron repulsion between both electrophores. A typical example is biphenyl, which is reduced to the radical anion at −2.68 V and to the dianion at −3.18 V (Figure 22.2) [10a,26], whereas the reduction of the monomeric benzene occurs at −3.42 V [28]. Obviously, the biphenyl anion has gained a considerable amount of conjugative stabilization energy, while in the dianion strong electron–electron repulsion dominates. On the other hand, the difference between the redox potentials of COT and bicyclooctatetraenyl is relatively small (−1.71 → −1.61 V vs. Ag/AgCl), indicating a low stabilization of the dimer in comparison to COT. Moreover, the Coulombic through-space repulsion between the excess charges in the dianion of the dimer is also small (ΔE0 ≤ 80 mV) [48,82]. However, it is unclear whether this dianion really represents a system with excess charges in the two COT units. It cannot be excluded that the first two redox steps generate the dianion of one COT subunit and only the trianion formation at −2.42 V (vs. Ag/AgCl) produces the first excess charge in the second COT subunit. In the course of the investigations, the concept of the oligomeric approach has been developed, which includes electrochemical studies of a great number of monodisperse chainlike hydrocarbons. Voltammetric measurements carried out with several oligomers of the p-phenylenevinylene (18, n = 1–6) [24] and the p-phenylene [26] series, respectively, clearly demonstrate that the reduction properties of such oligomers and polymers depend on the chain length of the systems. The following general trends have been developed as function of increasing chain length. First, the number of accessible redox states increases. The potentials of already existing redox states shift to less negative values when the next higher homologue is reduced. Obviously, the redox energies of different states gradually approach a common convergence limit with increasing chain length. Second, the redox states degenerate pairwise with increasing chain length, and third, in agreement with expectations, adding successive monomeric subunits in the molecular chain enlarges the stability of the system (Figures 22.2 and 22.3). From these results, it becomes clear that in charged oligomers a

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N=5 N=7

N = 10

N = 11

N = 15

N = 19

N = 23 –3.5 –3.0 –2.5 –2.0 –1.5 –1.0 –0.5

0.0

E (vs. Ag/AgCl) (V)

FIgURE 22.3 Voltammograms of the reduction of β-carotenoids (22) in DMA/0.1 M TBABr, T = 213 K, ν = 0.1 V/s, N = number of olefinic double bonds. The small waves in the reverse scans of N = 5 and 7 probably indicate dimeric coupling products of the anions or dianions. (With kind permission from Springer Science+Business Media: J. Solid State Electrochem., 2, 1998, 102, Heinze, J., Tschuncky, P., and Smie, A.)

reasonable number of energetically low-lying redox states are degenerated, followed by redox states with increasingly higher energies [10a]. π

π

=

18

π

=

19

π

=

20

π π n

Quite a large number of publications that have appeared since that time support these findings, but have also introduced new aspects that show the complexity of redox mechanisms in such systems [85]. Very systematic studies have been carried out by the Müllen group, who have varied in chainlike oligomers the type and coupling position of the electroactive monomeric building blocks and the modes of linkage, using both saturated and unsaturated species with different lengths [86].

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Their results clearly show that the size of the aromatic subunit, the overall number of π-electrons, the π-topology and steric effects are important factors for the redox behavior of all these species. Thus, in the series of oligoarylenevinylenes, replacement of the phenylene unit by larger arylene units enlarges the charge capacity of the respective systems. While the p-oligophenylenevinylene 18 (n = 1) with three phenylene units can be electrochemically reduced up to a dianion, the corresponding naphthalene derivative 19 (n = 1) reaches a trianion level and the anthracene derivative 20 (n = 1) even a hexaanion state [86–88]. An important reason for this phenomenon is the fact that the better the excess charges in condensed aromatic units are stabilized, the larger the π-structure is. Of course, the energetic stabilization of an excess charge in a large aromatic unit diminishes the trend for its delocalization and Coulombic repulsion effects along the chain. Therefore, the shift of the first redox potential in dependence on the chain length is less pronounced for the naphthalene and anthracene derivatives than for the phenylene system or the pure oligoene chain, and, moreover, the separation between successive redox steps becomes substantially smaller as the number of π-subunits increases. A further influence on the redox properties results from the coupling pattern between the vinylene and the arylene units. Measurements on para-, meta-, and ortho-coupled phenylenevinylenes reveal that it is more difficult to charge meta- and ortho-homologues than the corresponding para-homologues. However, the conjugative uncoupling of two meta-groups in a phenylene ring diminishes Coulombic repulsion, and therefore the energetic separation between successive redox steps decreases. The redox properties of oligo-p-phenylenes change when sterically relevant methyl groups exist in the central rings. Thus, in comparison with the unsubstituted oligomers in methyl-substituted homologues, for example, 21, with four or more phenylene units, the first reductive redox step is shifted to a more negative potential and a two-electron wave appears in the voltammetric response [89]. This can be interpreted by assuming that, due to steric hindrance, additional energy is needed to planarize the phenylene chain for the first electron transfer, and that the second electron is able to enter the then flattened system at the same or even a more positive potential.

21

Within the hydrocarbons containing olefinic double bonds, polyacetylene (PA) is the most thoroughly studied system. The great interest results from its extremely high conductivity up to 100,000 S/cm, which emerges on oxidative or reductive charging of the polymer [77–79,90]. As already found by MacDiarmid [9a], PA can be reversibly reduced to a polyanionic material. Despite the great interest in PA, there are only a few electrochemical studies of monodisperse oligoene systems. The reason is the high reactivity of doped alkyl-substituted oligoenes in the presence of nucleophiles or electrophiles. Normally, these oligomers consist of a carbon chain with alternating single and double bonds and two terminating groups, which are equal in most cases. Thus, a t-butyl group [91–93] or, in the case of α- and β-carotenoids, a cyclohexenyl group [94], has been used, while phenyl or other aromatic substituents have been used as end groups in the so-called arylpolyenes [95]. The chain length N (N = number of double bonds in the conjugated system) again determines the electronic properties of the oligomers.

22

n

n (N = 0–4)

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A series of carotenes [96] 22 (number of double bonds N = 5 and higher; compounds with N = 5 and 10 are not shown in structure 22) may serve as an illustration of this redox behavior. For all related states within the series, a strictly linear dependence of the redox potentials versus the chain length of the oligomers is observed. As can be seen, the reduction for the oligomer with N = 5 starts with two well-separated one-electron redox steps. With increasing chain length, there are additional weakly separated redox pairs. The potential gaps in the single pairs and between them decrease (Figure 22.3). Thus, two-electron transfer steps are most likely to occur for the longer polyenes (N ≥ 19). In the literature, even potential inversion for the second e-transfer is discussed [94]. The factors that control this inversion are essentially two. One is the weakening of the Coulombic repulsion brought about by the localization of the two charges of the di-ion at the end of the molecule. Additionally, localization of the charges at the end of the di-ion contributes to its stabilization by interaction with the solvent. In the case of the mono-ion, this stabilization is significantly lower. A further interesting point concerns unusually small differences between the reduction potentials of mono- and dianions of some molecules containing olefinic double bonds. Although in transstilbene the formation of the dianion occurs approximately 500 mV negatively to the radical anion 0 formation [26,97], for tetraphenylethylene the standard potentials for the R/R− couple ( ER/R − ) and 0 the R−/R2− couple ( ER − /R2− ) are very closely spaced. The ΔE 0 separation amounts in HMPA to −138 mV, in DMF to −35 mV, and reaches in ACN even positive values of approximately 150 mV [98–100]. Consequently, the disproportionation constant K varies within five orders of magnitude, a phenomenon mainly ascribed to increasing ion pairing on going from HMPA to ACN [100]. Nevertheless, intermolecular phenomena alone are not sufficient to explain the dramatic decrease of the ΔE 0 separation in tetraphenylethylenes in comparison to stilbene. It is now generally accepted that structural changes involving twisting of the ethylenic bond and accompanying the dianion formation are the main reason for the energetic stabilization of the dianion of tetraphenylethylene [101]. Similar effects also observed with 9,9′-bifluorenylidene [102].

23

24

25

Apart from investigations of chainlike conjugated systems, studies of molecules including orthogonal [103] as well as parallel π-systems [104] such as [2.2]paracyclophanes have attracted considerable interest. In that case, the intramolecular electronic interactions between the π-segments depend on the position of the connecting alkane bridges. Thus, the voltammetry of the dianion of 23 (∆E 0 = E10 − E20 = 420 mV) with a face-to-face arrangement of anthracene units indicates a strong through-space repulsion. This is lowered by going to the 1,4-bridged analog 24 (ΔE0 = 275 mV) or the orthocyclophane 25 (ΔE0 ≤ 80 mV) [104a]. Substitution of anthracene by pentacene electrophores weakens the repulsion [104b]. The reduction behavior of cyclooctatetraphenophane [56] shows one characteristic similarity to that of bicyclooctatetraenyl. Again, the ΔE 0 separation between the first two redox steps is small (≤80 mV), indicating a weak Coulombic repulsion between the negative excess charges in both COT subunits. In principle, one may also assume the formation of a dianion with two electrons in one

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subunit. Then the ΔE0 separation (0.38 V) between the redox potentials of the dianion and trianion would reflect the Coulombic repulsion between a dianionic and monoanionic subunit. Interestingly, the reduction of cyclooctatetraphenophane to the dianion occurs at potentials of about −2.50 V. Obviously, a considerable steric strain within the cyclooctatetraenyl subunit has caused this strong potential shift in comparison with the reduction potential of COT (−1.71 V). R

R 26

R

R

A very interesting group of [2.2]paracyclophanes are dibenzoannelated compounds of the general structure 26 with mutually orthogonal π-systems [103]. In the literature it has been speculated that the orthogonal π-systems of 26 are sufficiently isolated to be reduced independently. Cyclic voltammetric measurements show, however, that there are considerable interactions between the “isolated” π-systems. Thus, the unsubstituted [2.2]paracyclophane containing only two benzene groups in face-to-face position is reduced to its radical anion at −3.01 V [105]. By contrast, the benzoannealed cyclophane 26 with R = tert-butylphenyl could be reduced to its tetraanion (E10 = −2.54 V, E20 = −2.61 V, E30 = −2.80 V, E40 = −2.89 V). In conjunction with ESR measurements [106] and electrochemical data of ortho-terphenyl (E10 = −2.62 V, E20 = −2.72 V) [103], this gives evidence that the mono- and dianion formation of 26 takes place in the lateral subunits whereby the redox states are almost degenerated and easier accessible than in o-terphenyl. The most interesting finding is that the third electron is localized in the central paracyclophane subunit and its reduction potential is considerably less negative than that of [2.2]paracyclophane, given the fact that both annealed o-terphenyl units are already charged. A new class of conjugated hydrocarbons is that of fullerenes [9b], which represent an allotropic modification of graphite. Their electrochemistry has been studied in great detail during the last decade [107] (see Chapter 21). The basic entity within this series is the C60 molecule. Due to its high electron affinity, it can be reduced up to its hexaanion [10c,108]. Solid-state measurements indicate that the radical anion of C60 reversibly dimerizes. NMR measurements confirm a σ-bond formation between two radical anion moieties [109].

C.

ELECTRON TRANSFER KINETICS

Conversion of an oxidized species into the reduced form and vice versa requires the reorganization of the solvent in the immediate neighborhood of the reactant, together with some structural changes within the reactant. In the case of heterogeneous charge transfer, additional double layer effects are operative, which depend, inter alia, on type and concentration of the supporting electrolyte, but may be also influenced by the electrode material used as well as adsorption phenomena. Theoretical concepts have been developed by Hush, Marcus, and Dogonadze [110–112]. Applications of the Marcus theory to problems in organic electrochemistry have been discussed by Eberson [113]. As most studies on the reduction of hydrocarbons were carried out in aprotic solvents in the presence of excess supporting electrolyte, double layer influences on electron transfer kinetics were usually regarded as less important [2,114]. Normally, the rates of heterogeneous electron transfer to aromatic and olefinic hydrocarbons are high. It is assumed that the activation barrier is mainly caused by the solvent reorganization. Such reactions are termed outer-sphere processes. Their heterogeneous rate constants may reach values up to 5 cm/s [115]. Reductions of systems that require in addition a large conformational energy change in the transition state are rare [116]. The most widely discussed example from this class is the COT molecule (1), which exhibits a slow heterogeneous electron transfer for the anion formation, while the rate between the anion and the dianion is significantly faster [46,47,52].

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In the literature, it has been suggested that upon reduction the tube-shaped ring passes through at least a partially flattened transition state to form a planar anion radical. The experimentally determined activation energy (10–11 kcal/mol), which is similar to the activation energy of ring inversion (13.7 kcal/mol), has been interpreted as supporting evidence [114]. However, independent studies of Fry et al. [36] and Parker et al. [117] have shown that the rate constant for the first reduction step increased as the cation size of the supporting electrolyte decreased from tetraheptylammonium to tetramethylammonium. This contradicts the assumption of pure structural reorganization effects during the reduction of 1 and points to adsorption phenomena of the electrolyte cations. Nevertheless, in a very recent EPR investigation, homogeneous rate constants for the COT− •/COT electron self-exchange process were found to be close to 10 6 M−1 s−1, which is a thousand times slower than those of planar aromatics and confirms the concept of a strong inner-sphere reorganization upon electron transfer [118]. Similarly, Evans discussed electron transfer reactions of some fully α-methylated cycloalkane-1,2-diones in which the contributions of internal reorganization are substantial [119]. In that case, both homogeneous (self-exchange) and heterogeneous electron transfer rate constants are affected by the structural differences among these diketones. Tetraphenylethylene is another example of a slow charge transfer. Here, the radical anion formation is relatively fast, whereas the second charge transfer is slow [98]. In agreement with the thermodynamic findings, the second charge transfer suggests that considerable structural reorganization occurs at the activation barrier of the second reduction step. It is not yet clear to what extent solvent reorganization as well influences the activation energy [98].

III. CHEMICAL REACTIONS OF ELECTROgENERATED ANIONS With one or even more electrons in antibonding orbitals, reduced hydrocarbons are highly reactive species capable of both inter- and intramolecular reactions. The preferred pathway of the follow-up reaction depends not only on the electronic and steric structure of the reduced species but also on its chemical environment, especially on the counterion and the solvent.

A.

HOMOGENEOUS ELECTRON TRANSFER

The most elementary follow-up reaction is the homogeneous electron transfer (ET) to another solution species, which may be identical with the donor molecule itself [43a]. The equilibria of homogeneous ET reactions are governed by the standard potentials of the involved redox couples and are easily calculated with given data according to Equations 22.6 through 22.9: K

−•  ⇀ A1− • + A 2 ↽  A1 + A 2 E10

−•  ⇀ A1 + e − ↽  A1

E20

(22.6) (22.7)

−•  ⇀ A2 + e− ↽  A2

(22.8)

RT ln K = E20 − E10 F

(22.9)

The kinetics are much more complex and depend on the reorganization of the molecular framework [120], the solvation shell, and the electrostatic interaction. A semiquantitative estimation of rate constants may be obtained with the well-known Marcus equation [121]. The calculated data compare

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quite well with experimental values. Most of the experimental hydrocarbon data have been provided by Szwarc and his school [5]. The state of the art has been discussed in an excellent review [113]. Homogeneous ET has to be taken into account whenever dealing with electrode reactions, as was pointed out by Marcoux [122]. It is of special importance when the heterogeneous ET is slow, providing an additional and more effective pathway to reduced species. The +e 1 slow  → 1− •

−

(22.10a)

−

(22.10b)

e 1 + → 12−

↑-----------↓ 12 − + 1  2 × 1− •

(22.10c)

homogeneous rate constant for the second charge transfer in 1, similar to the heterogeneous charge transfer, is considerably larger than the first one. The difference greatly depends on the experimental conditions. It has been shown by cyclic voltammetry in superdry THF/TBAPF6 that the first ET can be catalyzed by the dianion of COT itself in a homogeneous synproportionation (Equation 22.10), giving rise to a catalytic spike (Figure 22.1). The mechanism has been confirmed by digital simulation [50]. Quite often the resulting anionic species undergoes a fast follow-up reaction. Thus, homogeneous ET becomes a crucial step in many electrode reactions: a well-known example is the cathodic reduction of organic halides (see Chapters 24 and 25). At mercury, heterogeneous electron transfer to most halides is slow. This kinetic barrier may be circumvented by addition of aromatic hydrocarbons “A” such as phenanthrene, naphthalene [123], and anthracene [124]. These compounds are easily reduced at the cathode and transfer their excess electron to the organic halide. The resulting halide anion radical undergoes a fast follow-up reaction. The whole reaction sequence is A + e− ⇌ A− •

(22.11)

A − • + PhCl ⇌ A + PhCl − •

(22.12)

PhCl − • → Ph • + Cl −

(22.13)

Ph • + A − • ⇌ Ph − + A

(22.14)

Ph − + H + → Ph H

(22.15)

Regeneration of hydrocarbon in Equations 22.12 and 22.14 makes the process catalytic. The reduction of organic halides in the presence of aromatic hydrocarbons has been the subject of detailed kinetic studies, which provide rate constants for the homogeneous ET [125] and the follow-up reaction [126]. The theoretical basis for this kind of experiment (homogeneous redox catalysis) was laid by Savéant’s group in a series of papers in 1978–1980 [127–129]. Homogeneous ET also plays an important role in the protonation of anion radicals [130]. When an anion radical undergoes heterogeneous ET, formation of the neutral molecule in the ground state is strongly favored over formation of an excited state [131]. No such restriction applies to homogeneous ET, which, if sufficiently exothermic, may yield excited states of hydrocarbons. One may naively suppose that an electron is removed from the bonding MO of highest energy to give either the first excited singlet or triplet; electrochemiluminescence [132] may then occur.

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The emission observed is usually the fluorescent band of the hydrocarbon, corresponding to the decay of the first excited singlet. The energy released by homogeneous electron transfer is given by the difference between the redox potentials of donor and acceptor plus a small entropy term [133], and in many cases of electrochemiluminescence falls short of the energy needed to populate the singlet state directly. In these energy-deficient cases, there is sufficient energy to populate the lowest triplet state, and singlets can then be produced by diffusion-controlled triplet–triplet annihilation. Emissions observed at wavelengths other than that of the main singlet have been ascribed to excited dimers (excimers) [134,135], excited charge-transfer complexes (exciplexes) [136,137], and fluorescent products of radical-ion decay. Acceptors in electrochemiluminescence may be the corresponding hydrocarbon radical cations, added alkyl halides [138] or benzoyl peroxide [139], or adventitious impurities, which need be present at only 10 −7 M levels. The experimental technique for the method is demanding [140].

B.

ELECTROPHILIC AND RELATED REACTIONS

1. Protonation Under protic conditions, aromatic hydrocarbons and compounds with activated double bonds usually undergo Birch-like reactions [141]. The reaction sequence has been elucidated by the classical work of Hoytink [11,12,142], who used HMO theory to rationalize both chemical and electrochemical steps. The anion radical produced by homogeneous ET is monoprotonated to give a Wheland-like π-radical. It readily accepts another electron because its bonding or nonbonding singly occupied molecular orbital (SOMO) always has a lower energy than the antibonding LUMO of the parent hydrocarbon. The second ET preferably occurs by disproportionation [130]. In a fast follow-up reaction, the resulting carbanion takes another proton to eventually yield the final dihydro product (Equations 22.16 through 22.20): A + e− ⇌ A− •

(22.16)

A − • + H + ⇌ AH •

(22.17)

AH • + e − ⇌ AH −

(22.18)

AH • + A − • ⇌ AH − + A

(22.19)

AH − + H + → AH 2

(22.20)

or

If the reduction potential of the resulting dihydro product is sufficiently positive, it can undergo another reduction cycle. One of many examples is provided by the reduction of benz[a]anthracene in 75% aqueous dioxane [11,12,143]. Because of its general importance to organic electrochemistry, the reaction scheme just outlined has been the subject of detailed mechanistic studies [16,144–151]. As a model reaction, protonation of anthracene anion radicals by phenol in dimethylformamide has been selected. Of five limiting kinetic variants, ECErev, ECEirr, DISP1, DISP2, and DISP3, the favored pathway was found to be DISP1 [130]. (The different types of rate-determining disproportionation reactions have been discussed in Reference 130.) It involves the protonation of the anion radical as the

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rate-determining step (Equation 22.17), followed by homogeneous ET between the anion radical and the protonated anion radical (Equation 22.19), yielding the monohydrogenated anion, which is itself rapidly and irreversibly protonated to the final dihydrogenated product. This mechanism has been disputed because of apparent deviations from predicted reaction orders [145–148]. Recently, it turned out that the inconsistencies are largely due to the formation of homoconjugated complexes between phenol and the phenolate anion [150]. Thus, DISP1 now seems to be generally accepted. The rate-determining step of the DISP1 mechanism, the protonation of the radical anion, largely depends on its electronic structure. As a guideline, LUMO energies of the parent hydrocarbon may be used [32]. The attempt to correlate rate constants with highest local charge densities failed [152]. Therefore, Eberson suggested the application of the Dewar–Zimmermann rules [153]. Under highly protic conditions, the major products of cathodic reductions of cyclic conjugated hydrocarbons are usually dihydro derivatives [46,154]. In 2-methoxyethanol, for example, naphthalene yields 1,4-dihydronaphthalene [154] and COT provides mainly 1,3,6-cyclooctatriene [46]. Under nominally aprotic conditions, 1,2-protonation dominates in naphthalene. Reduction of naphthalene in anhydrous acetonitrile containing tetraethylammonium p-toluenesulfonate yields 1,2-dihydronaphthalene, which is subsequently reduced to tetralin [155]. Similarly, reduction of 1 in anhydrous DMF gives 1,3,5-cyclooctatriene almost exclusively [46]. The formation of thermodynamically stable products is most probably due to base-catalyzed isomerization. An interesting situation arises from the reduction of CH-acidic hydrocarbons because these compounds can undergo self-protonation. Actually, a voltammetric investigation of 1,3- diphenyl2-methylindene and 4,5-methylenephenanthrene in DMF/TBAP or DMSO/TBAP clearly indicated a DISP1 mechanism, analogous to that described earlier [156a]. Similar results have been obtained for variously substituted indenes where the stoichiometry is in perfect agreement with a two-electron, two-proton reduction process involving 1/3 starting material under selfprotonation conditions, the remaining 2/3 acting as a proton donor [156b]. Generally, under self-protonation conditions the DISP1 pathway operates (Equations 22.21 through 22.24), in which the protonation reaction between the radical anion (AH− •) and the neutral species (AH) is the rate-determining step: AH + e − ⇌ AH − •

(22.21)

AH − • + AH ⇌ AH 2 • + A −

(22.22)

AH 2 • + AH − • ⇌ AH 2 − + AH

(22.23)

AH 2 − + AH → AH 3 + A −

(22.24)

On the other hand, reduction of fluorene [157,158] results in a homolytic cleavage of the CH bond (discussed in Section III.D). To study the stereochemistry of protonation reactions, substituted indenes have been cathodically reduced in DMF/TBAP with added water or phenol [156b,159,160]. In the presence of water, a formal anti addition of protons was observed, whereas addition of phenol led to the prevalent formation of products, formally deriving from syn protonation. Obviously, steric effects and/or acidities of proton donor and the dihydro product play an important role. Quite often, ion pairing causes a substantial positive shift of the reduction potential. For electrostatic reasons, the shift is especially large for the formation of higher valency ions. Therefore, with increasing interaction of the counterion, di- and polyanion formation becomes thermodynamically more favorable. Under these conditions, cathodic reduction immediately produces dianions

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via disproportionation and heterogeneous ET. Because of their high basicity, the dianions readily undergo protonation. Such a dianion mechanism was observed when tetraphenylethylene was reduced in acetonitrile/TEAP in the presence of water or alcohols. Kinetic measurements led to a mixed-order rate law, rationalized by the existence of a hydrogen-bonded complex between the ionpaired dianion and the proton donor [100]. The classical Hoijtink mechanism and the dianion mechanism have been observed at electrodes with a high hydrogen overvoltage, such as mercury. If mercury is replaced by platinum with its low hydrogen overvoltage, a radical pathway seems to be favored [161a], which is closely related to catalytic hydrogenations of hydrocarbons. Spectroelectrochemical experiments provided evidence for an additional hydride mechanism (Equations 22.25 through 22.27) [161b]: 2H + + 2e − ⇌ H 2ads

(22.25)

• H 2ads + e − ⇌ H − + H ads

(22.26)

A + H − ⇌ AH −

(22.27)

These results illustrate that the reaction sequence and the stereochemistry of cathodic hydrogenations are controlled by the solvent, electrolyte, proton donor, electrode material, etc. Thus, with increasing mechanistic knowledge, electrochemistry offers the chance to realize highly selective hydrogenations. The addition of protons certainly is the most common nucleophilic reaction of reduced hydrocarbon species because of the almost ubiquitous availability of proton donors, which may be wanted or not. However, under strictly aprotic conditions it is also possible to add other electrophilic reagents such as alkyl halides, acyl halides, CO2, and SO2. 2. Alkylation The addition of alkyl halides to aromatic anion radicals, generated by alkali metal reduction in ethereal solvents, was already known in the 1950s [162] and was reviewed by Garst in 1971 [163]. The first electrochemical analog was observed by Lund et  al. [164]. These authors cathodically reduced hydrocarbons such as naphthalene, anthracene, stilbene [123,124], and perylene [125] in the presence of alkyl halides and isolated hydrogenated and alkylated products. Similar reactions are observed when the halides are replaced by ammonium or sulfonium [165]. These alkylations can be looked upon as aliphatic nucleophilic substitutions, usually thought to proceed via SN1, SN2, or hybrids of these mechanisms. However, in recent years more and more evidence for a single-electron transfer (SET) mechanism, represented in Equations 22.28 through 22.31, was obtained, and it was suggested that SN2 and SET are just limiting cases of the same SET mechanism [166]. The SET pathway involves first a transfer of an electron from the nucleophile to the electrophile followed by bond formation, whereas the SN2 reaction involves a synchronous shift of a single electron and bond formation (Equations 22.32 and 22.33). In addition, the generated anions may be protonated (Equation 22.34). Because of these fundamental aspects, the mechanism of cathodically induced alkylations has been the subject of detailed studies [125,167]. In a stereochemical investigation, it was found that racemization is much more effective than inversion. This result was interpreted as evidence of competition between the SET pathway and the SN2 mechanism, with SET being the more important route [167]. The SET mechanism is represented in Equations 22.28 through 22.34:

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A + e− ⇌ A− •

(22.28)

A − • + BX → A + B• + X −

(22.29)

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A − • + B• → AB−

(22.30)

A − • + B• ⇌ A + B−

(22.31)

A − • + BX → AB• + X −

(22.32)

A − • + AB• ⇌ A + AB−

(22.33)

B− , AB− + H + → BH, ABH

(22.34)

The rate-determining step (Equation 22.29) of the SET mechanism consists of an electron transfer concerted with a cleavage of the carbon–halide bond in the alkyl halide and resulting in the generation of an alkyl radical. Numerous investigations have focused on the measurement of these rate constants [125]. As expected, the rate constant increases when the redox potential of the aromatic compound becomes more negative. The coupling step between alkyl radicals and aromatic anions is fast, with rate constants at the level of diffusion control. This indicates the lack of significant activation barriers, which consequently results in insensitivity to structural differences in the alkyl radical and in the aromatic radical anion [168]. Moreover, radical anions with very different redox potentials (∆EA0 = 0.9 V) couple with primary radicals with approximately the same rate constant. The competing SN2 mechanism (Equations 22.32 and 22.33) may be favored if the reacting species are not too sterically hindered, and the driving force for an electron transfer reaction is low. In general, the more positive the redox potential of the aromatic compound is or the poorer the alkylhalide is an electron acceptor, the more important the SN2 mechanism becomes [169]. 3. Acylation Formally related reactions are observed, when anthracene [170] or arylolefines [171] are reduced in the presence of carboxylic acid derivatives such as anhydrides, esters, amides, or nitriles. Under these conditions, mono- or diacylated compounds are obtained. It is interesting to note that the yield of acylated products largely depends on the counterion of the reduced hydrocarbon species. It is especially high when lithium is used, which is supposed to prevent hydrodimerization of the carboxylic acid by ion pair formation. In contrast to alkylation, acylation is assumed to prefer an SN2 mechanism. However, it is not clear if the radical anion or the dianion are the reactive species. The addition of nitriles is usually followed by hydrolysis of the resulting ketimines [171]. 4. Addition of CO2 In the pioneering publications of Wawzonek et al. [13,172], it was demonstrated that CO2 can react with cathodically reduced hydrocarbons to yield dihydrodicarbonylates. Examples of this kind of reaction described in the literature include naphthalene, phenanthrene, anthracene, and 9,10-diphenylanthracene. An ECE mechanism was proposed by several authors [172,173]. This includes the generation of the radical anion of the hydrocarbon, its nucleophilic addition to CO2, and a second electron transfer involving an additional coupling of CO2 with the dianion. However, studies of Lund and Simonet indicated that additional mechanistic variants such as redox catalysis should be considered [174]. A very careful analysis of possible reaction pathways was carried out by Avaca et al. [175]. They showed that the overall reaction mechanism could be different depending on the reduction potential of the aromatic hydrocarbons relative to that of CO2. There exist at least two essential variants. CO2 is reduced at more positive potential than naphthalene or phenanthrene. Thus, the first steps of the reaction are the reduction of CO2 to its radical anion and the subsequent coupling between

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Organic Electrochemistry

CO2− • and the neutral aromatic. Then, a further electron transfer occurs accompanied by a second coupling with CO2. In the case of anthracene or 9,10-diphenylanthracene with reduction potentials positive to that of CO2, a DISP1 mechanism has been confirmed using different electrochemical techniques. This is similar to that proposed for aromatic hydrocarbons in the presence of proton donors.

C. REDUCTIVE COUPLING Cathodically reduced hydrocarbons undergo not only homogeneous ET and nucleophilic attack but also coupling reactions resulting in hydrodimerization and polymerization. Reduction of stilbene [13] or diphenylacetylene [172] in DMF yields 1,2,3,4-tetraphenylbutane, whereas phenanthrene [172] provides 9,9′,10,10′-tetrahydro-9.9′-biphenanthrene. Hydrodimerization was also observed with benzalfluorene [176]. If DMF is replaced by acetonitrile, protonation completely dominates hydrodimerization [13]. In carefully dried ethers, using alkali or alkaline earth metals salts as supporting electrolyte, 1,1-diphenylethylene can be reduced cathodically to give stable solutions of 1,1,4,4-tetraphenylbutane dianions [177]. These dianions can be cleaved by flash photolysis in the presence of excess 1,1-diphenylethylene to give transient anion radicals of 1,1-diphenylethylene. Kinetic analysis of the subsequent recombination confirmed the postulated radical–radical mechanism (RR route); the rate constant was found to be 0.5 × 109 M−1 s−1 [178]. In the literature, the nucleophilic attack of an olefinic radical anion on the double bond of a second neutral olefin, the so-called RS route (radical–substrate coupling), has been postulated as a mechanistic variant for hydrodimerization reactions, but unambiguous experimental results have not been presented [179]. 9-Cyanoanthracene undergoes a reversible electrodimerization with almost no side reactions, making a perfect model compound [180–184]. Again, it turned out that the dimerization follows the RR route. It is also interesting to note that water accelerates the dimerization. This effect was rationalized by specific salvation [184]. A careful study of the dimerization kinetics of 9-cyanoanthracene in different solvents and at low temperatures gives evidence that the coupling reaction is diffusion controlled and that its rate constant increases with increasing polarity of the solvent as predicted by the Debye–Smoluchovsky theory [185]. It should be noted that several authors have suggested a more complex reaction pattern, which at least involves a two-step mechanism [186]. Many alkenes, activated by electron-withdrawing groups, readily undergo hydrodimerizations. The best known example is the electrodimerization of acrylonitrile, the base of the commercial Monsanto process [187]. Evidence is presented there that the essential step is a coupling of two radical anions (RR route). Using the SECM technique, Bard has also shown that only the radical anions of acrylonitrile dimerize [188]. Sterically less demanding arylalkenes and dienes undergo not only dimerization but also polymerization. Styrene is polymerized in ethers by alkali metal reduction [189] or addition of cumyl potassium [190]. The mechanism of ET-induced polymerization was extensively studied by Szwarc and his school [191]. It turned out that the first step is dimerization of the styrene anion radical, usually obtained by addition of sodium naphthalene. Under aprotic conditions, the resulting dianion adds to monomers, forming polymeric living anions. It is interesting to note that the rate of the polymerization largely depends on the counterion. With conventional techniques and electrolytes, it was not possible to obtain living anions because they are rapidly protonated by tetraalkylammonium salts and residual water. The first report of the production of living polymers by an electrolytic method has to be attributed to Yamazaki et al. [192], who used tetrahydrofuran as solvent, and LiAlH 4 or NaAl(C2H5)4 as electrolyte for the polymerization of α-methylstyrene. A similar technique was used to polymerize styrene as well as derivatives [193–195]. The suggested mechanism agrees with the pathway described earlier.

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Aliphatic and Aromatic Hydrocarbons

D.

881

INTRAMOLECULAR REACTIONS

For many years, intramolecular reactions such as conformational changes, bond cleavage, bond formation, and valence isomerizations have been observed only when hydrocarbons were reduced with alkali metals in ethereal solvents. In most electrochemical experiments, these reactions were dominated by the electrophilic processes already described here. However, progress in experimental techniques [8,22,23] has made these reactions accessible to electroanalytical investigations, providing new mechanistic insight. 1. Conformational Changes In recent years, it has become more and more obvious that ET is frequently accompanied by conformational changes. The interconversion may precede or follow ET; the majority of these processes have to be classified as CE (as yet this mechanism has not been observed in the hydrocarbon series), EC, EEC, or ECE mechanisms, whereas only a few systems follow an EE pathway. An example for an EC process is the interconversion of the anion radical of cis-stilbene, which is quite slow on the voltammetric time scale [97], whereas cis-azobenzene anion radicals isomerize very rapidly [196]. Tetraphenylethylene undergoes two closely separated reversible additions of one electron [99,197]. The small difference of standard potentials, equivalent to a high disproportionation constant (see Equation 22.5), has been interpreted as interconversion from an almost planar to an orthogonal conformation when going from the neutral molecule to the dianion [198]. The thermodynamics and kinetics of ET-induced interconversions of substituted tetraphenylethylenes have been studied in great detail [120]. Reduction of 1,6-dimethylbicyclo[4.4.1] undeca-2,4,7,9-tetraene appears to proceed via an ECE scheme in which the initial reduction gives an unstable radical anion, which undergoes a structural change, giving another anion with a conjugatively stabilized π-system [65]. Conformational changes concurrent with ET (EE pathway) are observed upon reduction of 1, as cited earlier. Details were reviewed by Evans and O‘Connell [120]. 2. bond-breaking and bond-Making Reactions Reductive cleavage of carbon–carbon bonds was already observed in the 1920s by Ziegler and Thielemann upon alkali metal reduction of diarylalkanes in ethereal solvents [199]. As was shown by Lagendijk and Szwarc [200] for 1,2-di(α-naphthyl)ethane, the primary anion radical undergoes homogeneous disproportionation, which is supported by ion pairing. The resulting dianion decomposes by the fission of the CH2–CH2 bonds into the salts of α-naphthylmethyl carbanions. Similarly, 9,9′-bianthryl can be cathodically cleaved into anthracene and 9,9′-dihydroanthracene plus small amounts of reduced dimers. The dianion mechanism is quite slow, whereas the tri- and tetraanions are supposed to decay rapidly [201]. ESR spectroscopical investigations of the anion radical of fluorene indicated a first-order decay, and it was concluded that the CH bond undergoes homolytic cleavage [202]. Voltammetric studies of fluorene in DMF reached the same conclusion [157]. The thermal and photochemical ring-opening reactions of cyclobutene are classical examples of pericyclic processes [203]. In 1976, Bauld et  al. described an ET-induced analog [204]. They observed that benzo- and phenanthrocyclobutene undergo ring opening upon alkali metal reduction and suggested an ECE pathway. A voltammetric study of cis- and trans-tetrahydro-1,2-diphenylcyclobutanephenanthrene in THF-NaBPh4 confirmed this mechanism. However, it turned out that the rate of the ring opening largely depends on the counterion. If ion pairing is prevented by addition of 15-crown-5, the reaction rate slows down dramatically. At −50°C, the anion radical of the Z isomer becomes stable on the voltammetric time scale, whereas the dianion exhibits a fast ring opening (Figure 22.5c) [205].

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882

Organic Electrochemistry R1 R2 X1

X1

– 2e + 2e

+ – X1,2

R2 R1

SCHEME 22.1

R1 R2

R1 R2 + – X1,2

+ 2e – 2e

X2

X2 R2 R1

R2 R1

Oxidation or reduction of 1,3-dimethylidenecyclobutanes.

Two-electron reduction [50] or oxidation [206] of 1,3-dimethylidenecyclobutanes yields bicyclo[1.1.0]butanes (Scheme 22.1). Cyclic voltammetry of 2,4-di-9H-fluoren-9-yliden-1,1,3,3-tetramethylcyclobutane in DMF at low temperatures has demonstrated that bond formation proceeds via an EEC mechanism; the rate constant has been found to be 20 s−1 [50]. The most interesting feature of the intramolecular bond formation is that it occurs at the dianion level (Figure 22.4). This proves the validity of the RR route even for intramolecular reactions despite a strong Coulombic repulsion. ET-induced cycloadditions of polycyclic olefins and cycloreversions of cyclobutane species have been studied by ESR spectroscopy [207]. Upon chemical and electrochemical reduction, 2,2′-distyrylbiphenyl rearranges by intramolecular coupling into a “bis-benzylic” dihydrophenanthrene dianion, which can be either protonated to a 9,10′-dibenzyl-9,10-dihydrophenanthrene or oxidatively coupled to a cyclobutane species. It is interesting to note that the intramolecular coupling between the styryl units takes place at the di- and triionic level, which is comparable with a radical–radical coupling (RR route, EEC mechanism) [205a]. The experiments are again evidence that RS coupling is generally improbable for ionic dimerization reactions. At low temperatures, the coupling rate between the negatively charged styryl units slows down and the

0.5 μA

10.0 μA

–1.0

–1.4

–1.8

–2.2

–2.6

E (vs. Ag/AgCl)/(V)

FIgURE 22.4 Cyclic voltammograms of the reduction of 2,4-di-9H-fluoren-9-yliden-1,1,3,3-tetramethylcyclobutane in DMF/0.1 M TMAPF6, T = 223 K, Pt electrode; (top) ν = 100 mV/s, (bottom) ν = 100 V/s.

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883

Aliphatic and Aromatic Hydrocarbons

+e –e –2.12 V

+e –e –2.21 V

–

– –

+e –e –2.81 V

Na+ (Li+)

hv

+e –e

+e –e

–

–1.35 V

A3–

+e –e –3.13 V

A4–

Na+ (Li+)

+e –e

– –

B3–

–2.98 V

Na+ (Li+)

hv

–

+e –e –2.60 V

2– +e –e 20°C, resulting in a curved Arrhenius plot and a deuterium kinetic isotope effect that decreased with decreasing temperature pointed toward a reaction scheme including the reversible dimerization of the radical cation followed by the irreversible formation of ArCH2+, ArCH3, and a proton (Equations 23.62 and 23.63): +

−d  ArCH 3•  dt +

2ArCH3

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+ 2

= k62  ArCH3•  +

(ArCH3 )2

(23.61) (23.62)

909

Oxidation of Hydrocarbons +

(23.63)

ArCH+2 + ArCH3 + BH+

(ArCH3 )2 + B

The implications of these unexpected results on the conclusions drawn from previous studies, in particular the work based on spectroelectrochemical measurements, are still not clear. The structure of the dimer dication (ArCH3+•)2, is not known. However, the reversible formation of dimers from hydrocarbon radical cations is not exceptional and has been observed in other cases [202–204] and is probably a general phenomenon, but usually the dimers do not manifest themselves kinetically. Another unexpected result is the observation that deprotonation of the radical cations of 9-methylanthracene and related substrates in the presence of pyridine bases appears to proceed via an initial complexation prior to the proton transfer (Equation 23.64) rather than by direct proton transfer (Equation 23.3) [205]. The important feature of the reaction is kinetic isotope effects that indicate a significant degree of quantum mechanical proton tunneling: CH3

CH3 + +

Base

+

Base

(23.64) CH2 +

BaseH+

Remotely related to this is the oxidation of, for example, 9,10-DMA, by Ce4+ in MeOH/MeCN that at 0°C leads exclusively to the 9,10-adduct, 19, which upon heating to 45°C was converted to the side chain oxidation product, 20, by elimination of MeOH [206]. CH3 OCH3

CH3

CH3 OCH3

CH2OCH3

19

20

In the absence of a proton in the α-position, for example, when the alkyl group is tert-butyl, the radical cation may instead undergo slow C–C cleavage resulting in formation of the parent hydrocarbon and a tert-butyl fragment. This has been observed during the voltammetric oxidation of 9-tert-butylanthracene in HFP [207].

D.

SIMPLE ALKENES, ALKADIENES, AND RELATED COMPOUNDS

The electrochemical oxidation of nonbenzenoid unsaturated hydrocarbons is believed to include also the formation of the corresponding radical cation as the first step. However, radical cations derived from alkenes are only in rare cases sufficiently stable to allow for the observation of their reduction back to starting material during CV. Accordingly, the oxidation potentials, such as those listed in Table 23.2, have in most cases no thermodynamic significance. A notable exception is the oxidation of hydrocarbons of the carotenoid type, which are oxidized in two closely spaced reversible or quasi-reversible one-electron transfers to the corresponding dications [189,190] and even the quasi-reversible oxidation to the radical trication has been observed [208].

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910

Organic Electrochemistry

E. ALKANES The anodic limit of MeCN containing a perchlorate salt is not high enough to allow a more general study of aliphatic hydrocarbons, except in a few favorable cases [174,209,210]. With the introduction of tetrafluoroborates and hexafluorophosphates [188] as supporting electrolytes the anodic limit in MeCN can be extended considerably, so that certain saturated aliphatic hydrocarbons can be studied voltammetrically. The application of ultramicroelectrodes has even made it possible to record voltammograms for short-chain alkanes in MeCN in the absence of intentionally added supporting electrolyte [211]. Oxidation potentials for a number of aliphatic and alicyclic hydrocarbons are listed in Table 23.3. The table is organized to include a representative range of alkanes; common conditions for voltammetry are displayed, as are the relationships between oxidation potentials measured against different reference electrodes. From the diffusion currents it can be estimated that two electrons are involved in the majority of these oxidations [188,209,212]. Radical cations of aliphatic hydrocarbons are not observed as a result of electrochemical oxidation, but insight into the nature of these species has been gained by high-level computational methods [213]. It will be apparent from Table 23.3 that, apart from the use of tetrafluoroborates and hexafluorophosphates, extended anodic limits can be obtained at low temperatures [215] or by the use of acidic

TAbLE 23.3 Oxidation Potentials of Alkanes E1/2 (V) Compound

vs. SCE

3.4a; 2.7b 3.5–3.9d 1.14–1.46e 1.7–2.3f 3.28a 3.01a

n-Alkanes

2,2-Dimethylbutane 2-Methylpentane Cyclopropane 1,1,2,2-Tetramethylcyclopropane Phenylcyclopropane Cyclopentane Cyclohexane Adamantane Cubane a b c d e f g h i j k l m

vs. Pd/H2

vs. Ag/Ag+ 3.5c

3.90c 1.68

g

3.41d 2.05h 1.87i 2.01g 1.77g 2.36j; 1.75k; 2.38l 1.73

m

MeCN–Et4NBF4 [188]. MeCN–Bu4NBF4 [214]. MeCN–Et4NBF4, −45°C [215]. MeCN [211]. V vs. Ag/AgCl, CF3SO3H,H2O [216]. HF–BF3/BF4− (Ho ≈ −15.6). FSO3H–CH3COOH [188]. MeCN–Et4NBF4 [106]. MeCN–LiClO4 [105]. MeCN–LiClO4 [217]. V vs. a Ag wire, CF3COOH–Bu4NBF4 [212]. MeCN–Bu4NBF4 [218]. MeCN–LiClO4.

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vs. Ag/Ag+

Ep (V) or Ep/2 (V)

Oxidation of Hydrocarbons

911

solvents such as CF3COOH [77,214,219], FSO3H [188,220–222], HF [223] and CF3SO3H (without supporting electrolyte) [216]. In addition to the electrolyte systems included in Tables 23.1–23.3, hydrocarbon oxidation has been realized in nitromethane, nitroethane, propylene carbonate, sulfolane, and dichloromethane [224]. The observed potentials, invariably for irreversible oxidation, vary in a reasonably predictable manner. Good correlations are often found with IP [210,212,225] and σ+ parameters [64,81,102]. There still remains some doubt about the first step of the overall reaction; however, in MeCN, the final products are usually the N-alkylacetamides, for example, as shown in Equation 23.36. In neutral solution, at extreme anodic potentials, it is difficult to decide between direct (Equations 23.1, 23.2, and/or 23.3) and indirect electron transfer (Equations 23.16, 23.17, and/or 23.18). For oxidation in MeCN–BF4− solutions, the variation in potentials is best explained in terms of the direct mechanism. An indirect oxidation mechanism involving hydrogen abstraction by electrogenerated nitrate radical has recently been proposed for the electrolysis of linear alkanes in tert-BuOH/H2O mixtures containing HNO3 and saturated with O2 [23]. In strongly acidic solution, it has been proposed that electron transfer is from the protonated alkane [222]. The CV of alkanes in FSO3H was studied as a function of added AcOH (a base under these conditions) and/or KFSO3; diffusion-controlled two-electron oxidations were found and the current–potential curves shifted as base was added. Similar results for cyclohexane oxidation in CF3COOH/FSO3H, CF3COOH/CH3SO3H, and CH2Cl2/FSO3H were also interpreted in terms of initial protonation of the alkane [219]. The experimental basis for this conclusion has been challenged. First, the addition of base-assisted alkane oxidation was claimed, which made it unlikely that oxidation of protonated alkane was involved [220]. Second, for cyclopentene, no oxidation peak could be observed for a solution of the alkene in FSO3H; 1H NMR spectroscopy showed that in such solution the cyclopentene was substantially converted by proton addition into the secondary carbenium ion [221]. It was suggested that this cation should be oxidized at a potential similar to those claimed for the oxidation of protonated alkanes. The issue has been confused further by the use of different reference electrodes and by variations in voltammograms with temperature that are difficult to interpret unambiguously as well as nonelectrochemical reactions of the alkanes in these highly acidic media [223]. However, there seems now to be general agreement on a mechanism involving electron transfer from unprotonated RH involving steps (23.1), (23.3), and (23.7).

V. CONCLUSIONS In most cases, hydrocarbon oxidations occur via an initial one-electron transfer to the anode to form a radical cation. The further fate of the radical cation depends on its reactivity toward other reagents present. By suitable blocking of reactive sites and/or extensive delocalization of the positive charge, the radical cation may be stable for long periods or at least during the time scale of slow sweep voltammetry. In the majority of cases, the radical cation is very reactive as an electrophile, proton donor, or electron acceptor and interacts with nucleophiles or bases present; cases are also known where it can react as a radical, that is, in dimerization processes. The factors that govern the reactivity of radical cations are becoming increasingly better understood as a result of kinetic work using conventional and voltammetric methods and also by theoretical studies.

REFERENCES 1. 2. 3. 4.

Fichter, F. Organische Elektrochemie; Steinkopff: Leipzig, Germany, 1942. Weinberg, N.L.; Weinberg, H.R. Chem. Rev. 1968, 68, 449. Lund, H. Acta Chem. Scand. 1957, 11, 1323. Sheldon, R.A.; Kochi, J.K. Metal-Catalyzed Oxidations of Organic Compounds; Academic Press: New York, 1981.

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Organic Electrochemistry

5. (a) Bard, A.J.; Ledwith, A.; Shine, H.J. Adv. Phys. Org. Chem. 1976, 13, 155; (b) Hammerich, O.; Parker, V.D. Adv. Phys. Org. Chem. 1984, 20, 55. 6. Courtneidge, J.L.; Davies, A.G. Acc. Chem. Res. 1987, 20, 90. 7. Pagni, R.M. Tetrahedron, 1984, 40, 4161. 8. Schmittel, M.; Burghart, A. Angew. Chem. Int. Ed. Eng. 1997, 36, 2551. 9. Yoshida, K. Electrooxidation in Organic Chemistry—The Role of Cation Radicals as Synthetic Intermediates; Wiley: New York, 1984. 10. Schmittel, M.; Ghorai, M.K. In Electron Transfer in Chemistry, Vol. 2; Balzani, V., ed.; Wiley-VCH: Weinheim, Germany, 2001, p. 5. 11. Hammerich, O.; Parker, V.D. J. Am. Chem. Soc. 1974, 96, 4289. 12. Dietz, R.; Larcombe, B.E. J. Chem. Soc. (B), 1970, 1369. 13. Hammerich, O.; Parker, V.D. J. Electroanal. Chem. 1972, 38, App. 9. 14. Wayner, D.D.M.; McPhee, D.J.; Griller, D. J. Am. Chem. Soc. 1988, 110, 132. 15. Marcoux, L.S. J. Am. Chem. Soc. 1971, 93, 537. 16. (a) Parker, V.D.; Eberson, L. J. Am. Chem. Soc. 1970, 92, 7488; (b) Parker, V.D. J. Electroanal. Chem. 1972, 36, App. 8. 17. Hammerich, O.; Parker, V.D. Electrochim. Acta 1973, 18, 537. 18. Bancroft, E.E.; Pemberton, J.E.; Blount, H.N. J. Phys. Chem. 1980, 84, 2557. 19. Aalstad, B.; Ronlán, A.; Parker, V.D. Acta Chem. Scand. 1982, B36, 199. 20. (a) Bard, A.J.; Phelps, J. J. Electroanal. Chem. 1970, 25, App. 2; (b) J. Electroanal. Chem. 1976, 68, 313. 21. Svanholm, U.; Ronlán, A.; Parker, V.D. J. Am. Chem. Soc. 1974, 96, 5108. 22. (a) Evans, D.H. Acta Chem. Scand. 1998, 52, 194; (b) Evans, D.H.; Hu, K. J. Chem. Soc. Faraday Trans. 1996, 92, 3983. (c) Evans, D.H. Chem. Rev. 2008, 108, 2113. 23. Tomat, R.; Rigo, A. J. Appl. Electrochem. 1986, 16, 8. 24. Ogibin, Yu.N.; Elinson, M.N.; Nikishin, G.I. Russ. Chem. Rev. 2009, 78, 89. 25. Wendt, H.; Schneider, H. J. Appl. Electrochem. 1986, 16, 134. 26. Chou, T.-C.; Cheng, C.-H. J. Appl. Electrochem. 1992, 22, 743. 27. Torii, S. Electroorganic Syntheses—Methods and Applications, Part I: Oxidations; Kodansha-VCH: Tokyo, Japan, 1985. 28. Shono, T. Electroorganic Synthesis; Academic Press: London, U.K., 1991. 29. Ogibin, Yu.N.; Nikishin, G.I. Russ. Chem. Rev. 2001, 70, 543. 30. Möller, K.-C.; Schäfer, H.J. Electrochim. Acta 1997, 42, 1971. 31. Ogibin, Y.N.; Ilovaisky, A.I.; Nikishin, G.I. Electrochim. Acta 1997 42, 1933. 32. Shono, T.; Nishiguchi, I.; Ohkawa, M. Chem. Lett. 1976, 573. 33. (a) Parker, V.D. Acta Chem. Scand. 1970, 24, 3151; (b) Parker, V.D. J. Chem. Soc., Chem. Comm. 1969, 848; (c) Parker, V.D. Acta Chem. Scand. 1970, 24, 3455. 34. (a) Barba, F.; Guirado, A.; Barba, I. J. Org. Chem. 1984, 49, 3022; (b) Barba, F.; Guirado, A.; Barba, I. J. Chem. Res. (S) 1986, 228; (c) Barba, I.; Gómez, C.; Chinchilla, R. J. Org. Chem. 1990, 55, 3272; (d) Barba, I.; Chinchilla, R.; Gómez, C. J. Org. Chem. 1991, 56, 3673; (e) Barba, I.; Tornero, M. Tetrahedron 1992, 48, 9967. 35. Barba, I.; Tornero, M. Tetrahedron 1997, 53, 8613. 36. Parker, V.D.; Dirlam, J.P.; Eberson, L. Acta Chem. Scand. 1971, 25, 341. 37. Shono, T.; Ikeda, A. Chem. Lett. 1976, 311. 38. Katz, M.; Saygin, Oe.; Wendt, H. Electrochim. Acta 1974, 19, 193. 39. (a) Kojima, M.; Sakuragi, H.; Tokumaru, K. Chem. Lett. 1981, 1707; (b) Inoue, T.; Tsutsumi, S. Bull. Chem. Soc. Jpn. 1965, 38, 661. 40. Ashikari, Y.; Nokami, T.; Yoshida, J. Org. Lett. 2012, 14, 938. 41. (a) Becker, J.Y. Israel J. Chem. 1985, 26, 196.; (b) Zinger, B.; Becker, J.Y. Electrochim. Acta 1980, 25, 791.; (c) Becker, J.Y.; Zinger, B. Tetrahedron 1982, 38, 1677.; (d) Becker, J.Y.; Zinger, B. J. Chem. Soc., Perkin Trans. II 1982, 395. 42. Manning, G.; Parker, V.D.; Adams, R.N. J. Am. Chem. Soc. 1969, 91, 4584. 43. Blount, H.N. J. Electroanal. Chem. 1973, 42, 271. 44. Evans, J.F.; Blount, H.N. J. Am. Chem. Soc. 1978, 100, 4191. 45. (a) Evans, J.F.; Blount, H.N. J. Electroanal. Chem. 1979, 102, 289; (b) Evans, J.F.; Blount, H.N. J. Phys. Chem. 1979, 83, 1970. 46. Svanholm, U.; Parker, V.D. Acta Chem. Scand. 1973, B27, 1454. 47. Shang, D.T.; Blount, H.N. J. Electroanal. Chem. 1974, 54, 305. 48. Schäfer, H.J.; Steckhan, E. Angew. Chem. Int. Ed. Eng. 1969, 8, 518.

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Oxidation of Hydrocarbons 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. 61. 62. 63. 64. 65. 66. 67. 68. 69. 70. 71. 72. 73. 74. 75. 76. 77. 78. 79. 80. 81. 82. 83. 84. 85. 86. 87. 88. 89. 90. 91. 92. 93. 94. 95. 96.

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914 97. 98. 99. 100. 101. 102. 103. 104. 105. 106. 107. 108. 109. 110. 111. 112. 113. 114. 115. 116. 117. 118. 119. 120. 121. 122. 123. 124. 125. 126. 127. 128. 129. 130. 131. 132. 133. 134. 135. 136. 137. 138. 139. 140. 141. 142.

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24

Activation of the Carbon– Halogen Bond Armando Gennaro, Abdirisak Ahmed Isse, and Patrizia Romana Mussini

CONTENTS I. II.

Introduction .......................................................................................................................... 917 Homogeneous Catalysis ........................................................................................................920 A. Metal Complexes ..........................................................................................................920 1. Cobalt and Nickel Complexes with Schiff Base Ligands ......................................920 2. Oxidative Addition to Palladium and Nickel Complexes ......................................921 3. Copper Complexes ................................................................................................. 922 B. Organic Mediators ........................................................................................................ 923 III. Heterogeneous Electrocatalysis ............................................................................................924 A. Problem of the Reference, Noncatalytic Cathode Material ..........................................924 B. Electrocatalytic Materials for C–X Bond Cleavage: An Overview ..............................924 C. Electrocatalytic Cleavage of the C–X Bond on Silver ..................................................926 1. Ag…X Specific Affinity: Primacy from Compromise ...........................................926 2. Specific Adsorption of Halide Ions on Silver Surfaces .........................................926 3. Role of DET Mechanism on the Electrocatalytic Effect of Silver......................... 927 4. Electrocatalytic Effects of Silver: Role of Surface Morphology ........................... 932 5. Electrocatalytic Effects of Silver: The Role of the Supporting Electrolyte........... 933 6. Electrocatalytic Effects of Silver: The Role of the Solvent ................................... 934 D. Palladium, Copper, Gold .............................................................................................. 935 References ...................................................................................................................................... 936

I. INTRODUCTION The injection of one electron into an organic halide, RX, by reaction with an electron donor in homogeneous conditions as well as by electrochemical or other means gives the fragmentation of the molecule, by breaking the carbon–halogen σ bond. For this reason, such a process is named dissociative electron transfer (DET). There are two possible reaction mechanisms for the reductive cleavage of carbon–halogen bonds (see Chapter 14). Electron transfer (ET) and bond breaking can occur either by a stepwise mechanism (Equations 24.1 and 24.2), with the intermediate formation of a radical anion RX.−, or in a concerted way in which ET and bond fragmentation occur in a single step (Equation 24.3): RX + e −  RX • −

(24.1)

RX • − → R • + X −

(24.2)

917

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918

Organic Electrochemistry

RX + e − → R • + X −

(24.3)

R• + e−  R −

(24.4)

RX + e–

Depending on the value of the applied potential with respect to the reduction potential of R∙, the  latter may undergo typical radical reactions or be reduced to a carbanion R− (Equation 24.4). In some cases, it is thus possible to trigger both a radical and an ionic chemistry [1]. In the case of the stepwise mechanism, two energy barriers must be taken into account: the first one is related to the ET (Equation 24.1), whereas the second is due to the carbon– halogen fragmentation (Equation 24.2), that is, the intramolecular ET from the orbital initially accommodating the incoming electron (generally a π* orbital) to the C–X σ* orbital (see Figure 24.1). A voltammetric investigation on an extended series of aryl halides, affording a regular, well-defined sequence of increasingly more stable radical anion intermediates,

RX

–

ΔG≠c ΔG≠e

Free energy

(a)

ΔG≠c ΔG≠e κ

(b)

ΔG≠c ΔG≠e (c) Reaction coordinate

FIgURE 24.1 Potential energy diagrams for a DET to RX under kinetic control of (a) the bond cleavage reaction, (b) both bond cleavage and ET reactions, and (c) ET. The subscripts e and c stand for the ET and bond rupture reactions, respectively. (Reprinted from Isse, A.A. et al., J. Phys. Chem. C, 113, 14983, 2009. With permission.)

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has shown [2] that a kinetic parameter κ, linked to the peak potential variation with scan rate in cyclic voltammetry (CV) ( RT /F ) (dEp /d log v)

(24.5)

RT  F ( Ep / 2 − Ep ) 

(24.6)

κ = –1.15 and to the CV half-peak potential width κ = 1.857

is an efficient indicator of the relative kinetic influence of the aforementioned two barriers. In particular, the relative importance of the first, electrochemical barrier with respect to the second one regularly decreases with increasing κ. When κ = 1, the DET process is kinetically controlled by the bond rupture (Equation 24.2), the ET step (Equation 24.1) being relatively fast. Conversely, when κ < 0.5, the ET step becomes the rate-determining step and κ coincides with the symmetry parameter of the energy barrier, α. The radical may undergo very interesting reactions. For instance, it may react with suitable nucleophiles present in the solution (Equation 24.7) to give nucleophilic substitution reactions [3] or it may dimerize (Equation 24.8) and/or disproportionate (Equation 24.9). A very important possibility is the reaction with suitable olefins (Equation 24.10), which is the key step in living radical polymerization processes such as atom transfer radical polymerization (ATRP) [4]: R • + Nu − → RNu• −

(24.7)

2R • → RR

(24.8)

2R • → RH + R(−H)

(24.9)

R • + C=C → R – C – C•

(24.10)

Also, the reactivity of the carbanion, for example, with suitable electrophiles, may be exploited for the development of important electrosynthetic pathways. In this connection, great interest is devoted to reactions with carbon dioxide (Equation 24.11) to give carboxylate compounds (see Chapter 25) [5]. Of course, the possibility of father–son nucleophilic substitution (Equation 24.12), which is a parasitic reaction leading to the dimer, must be considered and, if possible, avoided: R − + CO2 → RCOO −

(24.11)

R − + RX → RR + X −

(24.12)

The electrochemical reduction of organic halides is an important topic, not only for the stimulating and interesting mechanistic aspects, but also for the development of very useful electrosyntheses. In this connection, however, two major drawbacks must be considered. The first is the highly negative reduction potential required for the reduction of organic halides, which is an unfavorable aspect, both for the energetic cost and for the greater probability of concomitant undesired reduction processes. As is well known, the reduction potential becomes more negative in the series RI, RBr, RCl. The second problem is the possibility of parasitic reactions, in particular nucleophilic

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substitution, which lower selectivity and yields of the desired products. Furthermore, parasitic reactions are more important for the more easily reducible iodides, whereas chlorides are more resistant, but also more difficult to be reduced. For these reasons, much attention has been devoted to develop possible electrocatalytic processes for the activation of carbon–halogen bonds.

II. HOMOgENEOUS CATALySIS One line of action for the activation of carbon–halogen bonds is the homogeneous catalysis, which can involve an outer-sphere ET, between a suitable donor D.−, which can be produced by electrochemical reduction (Equation 24.13), and the halide (Equations 24.14 and 24.15): D + e −  D• −

(24.13)

D• − + RX → R • + X − + D

(24.14)

D• − + R • → R − + D

(24.15)

where the donor D.− can be either a metal complex or a relatively stable organic radical anion. However, in many cases, where the donor is a metal complex, Equation 24.14 involves an innersphere ET (ISET), with specific chemical interactions between the donor and the halide.

A.

METAL COMPLEXES

Several transition metal complexes, mainly with phosphine or nitrogen-based ligands, have been utilized for the activation of C–X bonds. The most widely investigated metals include Pd, Ni, Co, and Cu. The reaction mechanisms are strongly influenced not only by the nature and the chemical structure of the ligand but also by the chemical properties of the metal. 1. Cobalt and Nickel Complexes with Schiff base Ligands Cobalt complexes with tetradentate Schiff bases, such as salen (H2salen = N,N′-bis(salicylidene)-ethane1,2-diamine) and salophen (H2salophen = N,N′-bis(salicylidene)-phenylene-1,2-diamine) with different substituent groups, have been largely studied, in particular as model compounds of vitamin B12 [6]: O– O–

–O

–O

N N

N

N salen

salophen

In the case of Co complexes, CoII(L) is reduced to CoI(L)−, which reacts with RX to form an organocobalt complex RCoIII(L); this Co(III) complex is reducible at more negative potentials to give CoI(L)− and the radical R∙ [7–9]:

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Co II (L) + e −  Co I (L)−

(24.16)

RX + Co I (L)− → RCo III (L) + X −

(24.17)

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Activation of the Carbon–Halogen Bond

RCo III (L) + e −  RCo II (L)−

(24.18)

RCo II (L)− → Co I (L)− + R •

(24.19)

R • + e − [and/or Co I (L)− ]  R −

(24.20)

The reduction potential of RCoIII(L) is more negative than that of CoII(L), and it has been demonstrated that RCoII(L)− undergoes a homolytic fragmentation, even if ER0 i / R is more positive than the standard reduction potential of CoIII(L)R [9]. Reaction (24.17) has been shown to occur by an SN2 mechanism [10]. In contrast, a quite different behavior has been observed for Ni complexes with the same ligands. In this case, in fact, NiII(L) is reduced to NiI(L)−, which reacts with RX via an ET mechanism [11–14]: Ni II (L) + e −  Ni I (L)−

(24.21)

RX + Ni I (L)− → Ni II (L) + R • + X −

(24.22)

R • + Ni I (L)− → Ni II (L) + R −

(24.23)

2. Oxidative Addition to Palladium and Nickel Complexes Palladium complexes are largely employed as catalysts in several important reactions (see Chapter 36), in particular those involving C–C bond formation [15], such as arylation of alkenes [16,17] and alkynes [18,19], and cross-coupling reactions [20–26]. The key step of these reactions is the oxidative addition of aryl halides to a Pd(0) complex, as first reported by Fitton [27,28] for Pd0(PPh3)4 tetrakis(triphenylphosphine)palladium. This complex quantitatively dissociates to PPh3 and Pd0(PPh3)3 [29], but it has been demonstrated that the effective reactive species is Pd0(PPh3)2 in equilibrium with Pd0(PPh3)3 [30,31]: The kinetics of oxidative addition of aryl halides to Pd0(PPh3)2, quantified by k0a KL (Scheme 24.1), is quite insensitive to the solvent polarity [32]. This implies that the mechanism does not involve an ET as in the case of Ni0(PEt3)4 [33]. Pd0(PPh3)2 can be formed by electrochemical reduction of PdIICl2(PPh3)2. It has been demonstrated that the electrodic reduction of this precursor is really a bielectronic process that produces the anionic Pd0(PPh3)2X−, which is the best catalyst, since its rate constant for the activation of RX is higher than that of Pd0(PPh3)2 [34]. fast

Pd0(PPh3)4

KL

Pd0(PPh3)3 k0aKL [PPh3]

Pd0(PPh3)3 + PPh3 Pd0(PPh3)2 + PPh3 ArX k0a

= k app PPh3

X = I, Br

Ar

Pd

X

PPh3 Rate = k0a[ArX][Pd0L2] = k0aKL[ArX][Pd0L3]/[L] = k app[ArX][Pd0L3]

SCHEME 24.1 Mechanism of oxidative addition of ArX to Pd0 complex. (Reprinted from Jutand, A., Chem. Rev., 108, 2300, 2008. With permission.)

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Ni complexes can react by an ET mechanism [12–14,35–37] and also through oxidative addition [38,39] and SN2 [40,41] mechanisms, depending on the nature of the ligand. Square-planar Ni(I) complexes with Schiff base ligands [12–14] or tetraazamacrocycles [34–36] have been reported to react with alkyl halides by a DET (Equation 24.22). These complexes act as inner-sphere electron donors producing free alkyl radicals. Instead, nickel complexes with labile ligands such as halides, phosphines, and bipyridines can react through oxidative addition of RX. In the case of NiCl2(dppe) (dppe = PPh2–(CH2)2–PPh2), electrochemical reduction of Ni(II) produces Ni0(dppe), by two successive ETs, with the expulsion of the two chloride ions. Ni0(dppe) then reacts by oxidative addition with RX: Ni IICl 2 (dppe) + e − → Ni ICl(dppe) + Cl −

(24.24)

Ni ICl(dppe) + e − → Ni 0 (dppe) + Cl −

(24.25)

Ni 0 (dppe) + RX → RNi II X(dppe)

(24.26)

RNi II X(dppe) + e − → RNi I (dppe) + X −

(24.27)

In the presence of a suitable electrophile such as CO2, RNiI(dppe) reacts to produce Ni0(dppe) and RCO2− [38]. A similar reaction mechanism has been proposed also for nickel-bipyridine complexes [39]. 3. Copper Complexes Copper complexes have been largely investigated in the framework of controlled/living radical addition, cyclization, and polymerization reactions, in particular ATRP [4,42]. ATRP is, in fact, a controlled/living radical polymerization [43,44] used extensively for the preparation of homopolymers as well as random [45], gradient [46], block [47], graft [48], and dendritic polymers [49] with well-defined structures. It is based on a reversible halogen atom transfer (Scheme 24.2) between an alkyl halide RX (which can be either an initiator or a dormant macromolecular species Pm -X) and a low oxidation state metal complex (activator), resulting in the formation of propagating radicals (R• or Pm•) and the metal complex in a higher oxidation state (deactivator). N N

N N N N

N

PMDETA

Me6TREN

kp MtzLn + R-X KATRP =

kact kdeact

SCHEME 24.2

ATRP mechanism.

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kact kdeact

+ Monomer

X-Mt(z + 1)Ln + R kt

Bimolecular termination

923

Activation of the Carbon–Halogen Bond

Although several transition metals have shown catalytic properties toward various organic halides used as initiators [50,51], copper complexes with nitrogen ligands are the most used catalysts thanks to their low cost and easy handling; typically, they are prepared in situ by adding the ligand to a cuprous halide salt. The reaction between copper complex and RX has been shown to be an ISET, which proceeds with a reasonable rate even if the standard potential of the CuII/CuI couple is much more positive than that of RX. The ISET character of the reaction allows reduction of RX, but not that of R•, which otherwise would terminate propagation [52,53]. Metallic copper has also been used to activate alkyl halides toward formation of propagating radicals as well as regeneration of active Cu(I) catalysts in ATRP [54–56]. Knowledge of Cu speciation in ATRP conditions [57] allowed the study of the kinetics of activation of RX by a CuI complex in MeCN in both the absence and presence of halide ions. For the system CuI/L/X− (L = Me6TREN = tris(2-dimethylaminoethyl)amine), mainly composed of CuI(L)+, XCuI(L) and CuIX2−, only CuI(L)+ was found to be an active catalyst reacting with RX [58]. In recent years, the ATRP process has seen a significant progress by the development of a method known as activators regenerated by ET, which involves the use of reducing agents for the reduction of air-stable deactivators to their respective activators in solution [59]. These systems are conducted in the presence of excess reducing agent whereby the CuI/L/X− activator is continuously regenerated from CuIIX2/L, which accumulates as a by-product of unavoidable termination events. More recently, a very innovative ATRP process (eATRP) based on electrochemical generation/regeneration of the activator complex has been developed [60–62]. In this process, air-stable CuIIBr2/ Me6TREN is electrochemically reduced to CuIBr/Me6TREN to invoke or trigger polymerization. The feasibility of an electrochemical switch to modulate copper oxidation states in situ, and thereby activate or deactivate polymerization, was demonstrated.

B. ORGANIC MEDIATORS Aromatic radical anions (A• –) can act as outer-sphere electron donors toward organic halides [63]. In the case of an alkyl halide, RX, the principal steps of the reaction can be written as follows: A + e −  A• −

(24.28)

A • − + RX → R • + X − + A

(24.29)

A• − + R• → R − + A

(24.30)

A• − + R• → A − R −

(24.31)

This process is commonly known as homogeneous redox catalysis. It involves ET to RX, which undergoes C–X bond rupture, followed by a second ET to the ensuing radical R• and/or radical– radical coupling between A• – and R•. Some advantages of this method of C–X activation over metal catalysis include the ease of generation, for example, by electroreduction, of A• – from stable aromatic or heteroaromatic compounds; the low propensity of A• – to bond formation owing to the delocalization of the odd electron; the well-defined low self-exchange reorganization energy of the A/A• – couple [64], which makes easy analysis of the dynamics of reaction (24.29); and the nonspecificity of the reaction for different classes of RX. Activation of C–X bonds by electrogenerated aromatic radical anions has been used in different synthetic applications [13,65–67]. The homogeneous redox catalysis approach has also been widely used for measuring rate constants and theoretical analysis of DET processes [68]. Another important application of the homogeneous reduction of alkyl halides by A• – is the determination of

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Organic Electrochemistry

the reduction potentials of short-lived alkyl radicals, according to a method developed by Lund and coworkers [69,70] on the basis of the competition between reactions (24.30) and (24.31). This method has been successfully used for the estimation of standard reduction potentials of a large variety of alkyl radicals [71–74].

III. HETEROgENEOUS ELECTROCATALySIS A.

PROBLEM OF THE REFERENCE, NONCATALYTIC CATHODE MATERIAL

In the last decade, investigations on the electrocatalytic cleavage of carbon–halogen bonds have usually dealt with electrocatalytic effects in terms of differences between reduction peak potentials obtained in the same conditions on the electrode surface tested as a catalytic material and on a suitable electrode, acting as an inert (or noncatalytic) electron source ensuring outer-sphere ET. The catalytic effect is therefore defined as Catalytic effect = Ep,catalytic − Ep,noncatalytic

(24.32)

This unofficial convention has been proposed to provide a sort of normalization with respect to the substrate intrinsic reactivity, which is widely modulated by the molecular structure [75]. Glassy carbon (GC), boron-doped diamond (BDD), and fluorinated boron-doped diamond (FBDD) have been tested for this role [76]. The reduction potential sequence of different model organic halides on the three carbon-based electrodes is GC > BDD > FBDD, actually a sequence of decreasing Lewis acidity of the reacting surfaces, considering that the GC surface is partially functionalized with hydroxyl and carboxyl groups [77], while BDD is predominantly hydrogenterminated [78], and FBDD features a consistent number of fluoride atoms on its surface, resulting in an even more hydrophobic character [79], which however could exert a repulsive effect on the approaching halide leaving group of the reacting molecule. Accordingly, BDD would appear the most appropriate inert reference among the three surfaces tested. However, it is expensive and difficultly available, particularly in an electrode setup suitable for working in nonaqueous solvents; moreover, fundamental studies on organohalide electrochemical reduction in noncatalytic conditions have mostly been carried out on the more popular GC electrode; thus, the latter has been commonly adopted as the reference for evaluating catalytic effects in the same process. Care should, however, be exercised in the choice of GC electrodes when an accurate analysis of DET dynamics is desired since their electrochemical performance is strongly influenced by their surface microstructure [80,81]. In particular, edge-plane sites and defects are prone to oxidation (also depending on the adopted method of surface preparation [77]), and therefore, several oxygencontaining functional groups, mainly carbonyl, hydroxyl, and carboxyl moieties, are present on the electrode surface. This significantly affects the reduction peak potential of RX, an increase of the oxygen-to-carbon ratio apparently causing a positive shift of Ep [82]. This catalytic effect reproducibly depends on both GC electrode surface composition and DET mechanism, regularly increasing with increasing importance of the initial ET step in the overall kinetics of the process [82]. The highest observed peak potential difference between differently activated GC surfaces is on the order of 0.1 V, found for some alkyl halides undergoing a concerted DET [82].

B.

ELECTROCATALYTIC MATERIALS FOR C–X BOND CLEAVAGE: AN OVERVIEW

The use of electrocatalytic electrode surfaces is a convenient approach to the electrochemical activation of carbon–halogen bonds in mild conditions. All metals having specific halide affinity could potentially possess electrocatalytic activity for C–X bond cleavage; comparative investigations [76,83–89] encompassing a wide range of metals also in terms of electrocatalytic scale [76] or volcano plot [83] are available in the literature.

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Activation of the Carbon–Halogen Bond

925

The earliest studies on the electrocatalytic reduction of organic halides were rather involuntary (sometimes even unconscious), mostly consisting in organohalide reactivity studies using polarography, a ubiquitous voltammetric technique in the 1970s–1980s, implying to work on the noninnocent mercury electrode; actually halide adsorption phenomena and intermediate formation of organomercurials were found to play an important role in the process (see Chapter 25). Also, the first approach to silver in the 1970s was fortuitous, consisting in clinical observation of the anomalously high response of an oxygen Clark-type sensor in the presence of anesthetic halothane [90], an observation slowly exploited and rationalized in the following two decades [91–94]. Industrial applicative screenings also involved some early testing of catalytic properties of materials such as Pd and Ag [95], with some patents being deposited in the 1980s [96–100]. From the 1990s, synthetic [5,89,101–112] and, above all, environmental issues [87,88,113–126] have prompted remarkable acceleration, rationalization, and refinement in this research field. In particular, Hg was put aside because of environmental concerns and operating problems together with its only moderate catalytic effects [76,83], and investigations were concentrated on a series of solid electrodes that are significantly more catalytic than Hg, in particular Ag, which was found to possess a very high, reproducible and rationalizable catalytic activity for a very wide range of organic chlorides, bromides, and iodides [75,127,128], including mono- and polyhaloalkanes [75,126,129–132] as well as aryl [2,125,128,129,133], benzyl [1,134,135], glycosyl [106], and heteroaryl [108,128,136] halides. Silver nanoparticles [118,137] and alloys [115] have also been considered. Pd, both as such and as palladiated (even with nanoparticle Pd) surfaces of different nonmetals and metals, like GC, Ag, Cu, Ni, Pt, and Au [138–155], and Cu [86,89,121–123,156–158], have also been found to possess good electrocatalytic properties for the reduction of many classes of organic halides. Au is also sometimes considered [89,124,136,159–162]. It is important to stress, however, that although the catalytic activity of this metal for C–X cleavage is intrinsically higher than that of Ag [163,164], it is practically hampered by the huge negative charge density of the metal at the operating potentials (as a consequence of its more positive pzc) [165] unless some auxiliary condition prevails, for example, the availability of protons [159] or the presence of anchoring groups [136]. In recent years, the increasing interest of the process prompted detailed mechanistic investigations, particularly in the case of silver, on account of its remarkable and reproducible catalytic activity and of the availability of a wide series of authoritative studies on specific halide adsorption. Such investigations focused on different factors capable of modulating the process as, for instance, the reactant molecular structure (organic moiety and halide leaving group) [2,128,129], morphology [166], and state of cleanliness [127] of the electrode surface, adsorption phenomena involving the reagents and/or products [2,128,134,136,167–169], reaction medium (solvent and electrolyte) [120,135,159,170,171], and the presence of adsorption auxiliaries [136]. Exhaustive mechanistic rationalizations have been so far achieved for the electrocatalytic reduction of aryl [2] and benzyl halides [1,134,135,172,173], in the latter case also with the support of a combined computational and SERS investigation [172,173]. From the preparative point of view, the milder electrode potential and the involvement of the catalytic metal surface in the formation of reaction intermediates can promote intermolecular radical reaction pathways, affording even difficult dimerizations [1,105,106,110]. This radical chemistry also affords addition to suitable coadsorbed intermediates, as in the case of α-C- [101] and O-glycoside [102–104] synthesis from the catalytic reduction of α-acetobromoglucose. It is interesting to note that the same process at noncatalytic electrodes evolves through intramolecular reaction pathways hinging on a carbanion intermediate [174]. Instead, aryl halide reductions usually result in simple hydrogenation of the halide position [113]. The electrocatalytic properties of silver have already been applied (a) in several electrosynthetic processes, in particular the electrosynthesis of fine chemicals and pharmaceutical products by electrocarboxylation of the corresponding halides [5,107,109,111,124] achieving, in some cases, a first scale-up with very encouraging results [112], (b) for the treatment of environmentally relevant organic halide pollutants [87,88,113–121,125,126,155], and (c) for organic halide monitoring [85].

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C.

Organic Electrochemistry

ELECTROCATALYTIC CLEAVAGE OF THE C–X BOND ON SILVER

1. Ag…X Specific Affinity: Primacy from Compromise Quantum calculations in the vacuum on halide adsorption on the group 11 metals Cu, Ag, and Au point out that Au has the strongest affinity for X−, the general trend of halide affinity for all metals being I− < Br − < Cl− < F−, which is in disagreement with experimental results in solution [163]. DFT studies on halide interactions in the vacuum with three much more different metals, namely, Hg, Ag, and Pt, result in a metal sequence consistent with the experimental one (Ag > Hg > Pt) but, again, in the I− < Br − < Cl− < F− reversed halogen sequence; the disagreement with the experimental evidence is convincingly justified by the authors in terms of solvation energy contributions [164]. Analysis of halide-specific adsorption on metals in water in terms of competition between metal– water interactions (following the sequence Bi, Hg, Pb < Au(111) < Cd < Ag(111) < Ga < Cu(111)) and metal–halide interactions (following the sequence Bi, Pb < Hg < Cd, Ga < Cu(111) < Ag(111) < Au(111)) shows that the degree of surface coverage by X− replacing water molecules at the pzc increases in the order [175,176] Bi, Pb < Hg < Cd, Ga < Cu(111) < Ag(111) < Au(111) However, the catalytic activity of Au toward RX reduction is in practice greatly impaired [165] by its pzc (−0.05 V vs. SCE in NaClO4 aqueous solution [177] to be compared with −0.716 V vs. NHE [178] or −0.96 V vs. SCE for Ag), which is much more positive than the typical reduction potentials of the C–X bond. Although specific halide adsorption can take place even at potentials negative to the pzc, that is, with a negatively charged surface [179], increasingly negative charge densities do hamper specific interactions. Now, in the reduction potential range of most organic halides, silver is remarkably less negatively charged than gold. Accordingly, Ag is much more electrocatalytic than Au for this process in its usual working conditions. 2. Specific Adsorption of Halide Ions on Silver Surfaces Detailed studies on specific halide anion adsorption onto silver are available, most of them concerning investigations on monocrystalline Ag in water based on capacitance and zero-charge potential measurements [180–186], theoretical computations [187], computer simulations [188], and various other techniques, including impedance [189], chronocoulometry [190,191], electrochemical STM [192], surface X-ray scattering (SXS) [193], and surface electron spectroscopies (LEED, RHEED, and AES) [194]. These investigations show that specific adsorption of inorganic halides onto silver is so strong as to hold until threshold potentials negative with respect to the surface pzc, that is, about −0.8 V (SCE) for chlorides, −1.0 V (SCE) for bromides, and −1.2 V (SCE) for iodides. Moreover, the aforementioned studies very finely evidence a sequence of progressive structural transitions of the adsorbed halide anion monolayers with increasingly negative potential [190–192]. Recently, such investigations have been extended to controlled-surface polycrystalline silver [166–168]. The adsorption CV patterns observed on polycrystalline Ag, obtained by an appropriate electrodeposition protocol [169], are reproducible and intermediate with respect to the adsorption characteristics of the single crystals. Capacitive experiments showed the polycrystalline pzc to nearly coincide with that of the more open monocrystalline surface (110) [169]. More recently, such studies were extended to nonaqueous solvents such as acetonitrile [166,195], propylene carbonate [195], and dimethylformamide [195]), combining differential capacity and impedance experiments with an indirect voltammetric method based on the monitoring of the negative shift of the reduction peak potential of a “probe” organic halide molecule induced by progressive additions of halide anions, resulting in increasing adsorption competition [166]. Chloride, bromide, and iodide ions are

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Activation of the Carbon–Halogen Bond

specifically adsorbed onto polycrystalline silver electrodes in the three organic solvents studied, in the same Cl− < Br− < I− sequence as in water; threshold potentials for X− desorption in acetonitrile are ∼−1.20, −1.32, −1.43 V versus SCE, respectively [166]. The adsorption strength increases with decreasing solvent ability to coordinate the halide anions [185]. 3. Role of DET Mechanism on the Electrocatalytic Effect of Silver The electrocatalytic activity of silver toward the reduction of organic halides is strongly affected by the DET mechanism [2,86,128,129]. In the case of concerted DET, which is usually followed by aliphatic and benzyl halides, remarkable electrocatalysis is always displayed, independently of the nature of the halogen atom. It can be assumed [86] that an activated complex characterized by strong interactions between the reacting C–X bond and the electrode surface is formed; as a result, the intrinsic activation energy barrier (∆G0≠ = (λ 0 + BDE)/4, with BDE = bond dissociation energy and λ 0 = solvent reorganization energy) is significantly lowered. This decrease of ΔG ≠ results in a decrease in the required overpotential for the process and hence in a positive shift of E p with respect to the value observed on a noncatalytic electrode such as GC. The extent of catalysis in concerted DETs is also strongly affected by possible specific adsorption of the electrogenerated halide ion and/or the starting organic halide [134] if the reduction potential at Ag is not so negative that electrostatic repulsions prevent X− adsorption (see Section III.C.2). This specific adsorption effect decreases in the sequence I− > Br − > Cl−. Adsorption processes also result in alterations, sometimes conspicuous, of the reduction peak morphology with respect to the canonical diffusive feature. The case of benzyl halides PhCH2X provides a good example of the implications of reactant and product-specific adsorption as a function of the reduction potential (Figure 24.2). In the chloride case, the reduction potential is far beyond the negative threshold for Cl− adsorption; 0

PhCH2Cl

–20 –40

0.2 V/s

–60

I (μA)

0

PhCH2Br

–20 0.1 V/s

–40 0

PhCH2I

–10 –20 –30

0.1 V/s –2.0

–1.5

–1.0

–0.5

E (V vs. SCE)

FIgURE 24.2 CV patterns for the reduction of benzyl halides on Ag in MeCN. (Adapted from Isse, A.A. et al., Electrochim. Acta, 51, 4956, 2006. With permission.)

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Organic Electrochemistry

consistently, the molecule exhibits on Ag a single diffusive two-electron reduction peak (indicating that the radical undergoes easy reduction to anion at the same potential at which it is formed), although at a considerably less negative potential than at GC. A closer examination reveals that some slow adsorption of the reactant PhCH2X molecule does take place, although so weak that dissolved and adsorbed reactant molecules are reduced virtually at the same potential [134]. Very recently, a combined CV, SERS, and DFT investigation [172,173] has provided some insight into the process, confirming preadsorption of benzyl chloride, which after ET evolves into a benzyl radical bound onto the Ag surface; the latter is then easily reduced into a benzyl-silver anionic adduct that eventually dissociates. In the case of PhCH2Br, the DET process takes place within the potential range for specific Br− adsorption. At slow scan rates, the molecules exhibit two irreversible reduction peaks attributed to the reduction of PhCH2Br (first peak) and PhCH2•. The sharpness of the first peak indicates specific adsorption of one or both the reaction products. This peak gradually splits into two separate peaks as scan rate is increased. Addition of Br− in the reaction medium also brings about the same peak splitting even at low scan rates. These observations have been rationalized by considering the rate and degree of adsorption of Br− during the voltammetric scan [134]. In the iodide case, the DET takes place well before the threshold potential for iodide anion adsorption, and the adsorption of the PhCH2I reactant molecule at the Ag surface is so strong that reduction of adsorbed PhCH2I is observed as a sharp post peak at a potential about 150 mV more negative with respect to the diffusive one (Figure 24.2). The combination of all the aforementioned effects results in a scale of catalytic effects (anodic shift of Ep: 0.45, 0.72, and 0.48 V for X = Cl, Br, and I, respectively), which is at variance with the familiar silver/halide ion interaction sequence in solution, that is, I− > Br− > Cl−. It is also worthwhile noticing that in both cases in which significant halide anion adsorption can take place, that is, Br and I, a neat separation is observed between the potentials of the first and second electron uptake; this allows triggering selectively either radical or carbanion chemistry, as confirmed by preparative experiments [1]. The stepwise DET mechanism typically applies to aryl halides as a consequence of their ability to stabilize the incoming negative charge in a low-lying π* orbital delocalized on the aromatic system, so the ensuing radical anion energy curve intersects those of the reagent and product. This charge delocalization implies a much lower polarization degree and a higher BDE of the carbon halide bond to be cleaved with respect to concerted DET cases. Aromatic bromides show catalytic effects on Ag possibly arising from interactions between the electrode and various species, mainly ArBr • − and Br−, involved in the DET. These catalytic effects have recently been shown [2] to be regularly linked to the degree of negative charge localization on the halide atom in ArBr • −, suggesting that charge localization promotes specific interaction of the intermediate with the silver surface. Thus, for example, in the limiting case represented by 4-nitro-halobenzene, where the ET is circumscribed to the nitro group and therefore the negative charge in the radical anion is localized far away from the C–X bond, no electrocatalysis is displayed by Ag, independently of the type of the halogen atom [128]. Total absence of catalytic effects was also observed in the reduction of all aryl chlorides, at least in aprotic organic solvents [128,129]. This can be justified in terms of both the high value of the C–Cl BDE in aryl chlorides (396 kJ mol−1 for chlorobenzene) [196] and the negligible polarization of the C–Cl bond in the aryl radical anion intermediate ArX• − [197]. Therefore, with organic chlorides in aprotic solvents, electrocatalysis appears to be clearly discriminated by the DET mechanism, being remarkable in all concerted DET cases and absent in all stepwise ones [129]. As compared to aryl chlorides, aryl bromides have lower BDEs and give rise to radical anions with higher C–X bond polarization [197]. Accordingly, when dealing with a bromide leaving group, significant electrocatalytic effects of Ag are observed, regardless of the mechanism of DET. Such effects regularly depend on the molecular structure, and a rationalization has been recently achieved [2] by considering a systematic aryl bromide series with different substituents

© 2016 by Taylor & Francis Group, LLC

929

Activation of the Carbon–Halogen Bond Br R

Br 9 1

R=H C2H5

2

OCH3

3

C6H5

4

COC6H5

5

COCH3

6

CN

7

CO2C2H5

8

Br

10

Br

11

appropriately selected so as to modulate the kinetics of decomposition of the intermediate radical anion ArBr • −, from the limiting case of a process controlled by the first heterogeneous ET, as in the case of bromobenzene, to the other extreme, in which the process is controlled by the intramolecular ET, that is, the dehalogenation reaction, as in the case of bromoanthracene. These different kinetic regimes can be well described by a kinetic parameter κ (see Section I). As shown in Figure 24.3a, a plot of κ as a function of E p exhibits for both Ag and GC a linear increase of κ with increasing peak potential, which is increasing the electron-withdrawing ability of the substituent, or the number of fused aromatic rings stabilizes ArBr • − and consequently shifts the kinetic regime of the DET process toward kinetic control by the bond rupture reaction. We can see, however, a quite different slope for the two plots, which reflects a remarkable difference in the electrocatalytic effect played by Ag, which, in turn, is well correlated to κ (Figure 24.3b). The correlation between inductive effects in the aryl bromide molecular structure and Ag electrocatalytic effects is evidenced in the plots of Ep values at both GC and Ag electrodes versus Hammett substituent constants reported in Figure 24.4. As can be seen, both electrodes show the same trend; that is, Ep becomes more positive upon increasing the electron-withdrawing power of the substituent. For both electrodes, the dependence of Ep on σ− can be fit to a linear equation of the classical form. It is worth noting that reduction of the aromatic bromides at GC is ca. four times more sensitive to the substituents with respect to Ag since at the catalytic Ag electrode the

κ

4 0.6 2

3

7

6 75

8

8

1

1 3 2

0.2 –2.4 (a)

9 11

10

–2.0 Ep (V) vs. SCE

9 4

11 10

EpAg – EpGC (V )

6

0.8

0.4

1.0

5

1.0

0.8

2

3

1

0.6 0.4

11 4 9

0.0 0.4

–1.6 (b)

10

7

0.2

6 0.6

0.8

5 1.0

κ

FIgURE 24.3 Correlation between the kinetic parameter κ and (a) the reduction peak potentials on noncatalytic GC (squares) and catalytic Ag (circles) (Reprinted from Isse, A.A. et al., J. Phys. Chem. C, 113, 14983, 2009. With permission.) and (b) the catalytic effects of Ag, for the aryl bromide series.

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Organic Electrochemistry

Ep (V) vs. SCE

–1.6

3 2

5 6 6

1 –2.0

7

7

8 4

–2.4

3 2 1

–2.8 2 GC E Ag p – E p (V)

5

8

4

0.8

1

3 4

0.4

8

7

5

0.0 0.0

0.4

6

0.8

σ–

FIgURE 24.4 Reduction peak potentials measured on noncatalytic GC (circles) and catalytic Ag (triangles) and catalytic effects of Ag (squares) plotted versus σ− Hammett parameters, for the aryl bromide series. (Reprinted from Isse, A.A. et al., J. Phys. Chem. C, 113, 14983, 2009. With permission.)

molecular structure has a quite limited effect on Ep. It appears that the presence of the catalytic surface shifts, at least partially, the site of the negative charge in the 1e− reduced species. Probably now the charge is mainly located at the bromine atom so that the electron-withdrawing power of the substituent becomes less important. Consistent with such observations, a reduction mechanism has been proposed in which, as illustrated in Scheme 24.3, all reactions occur on the electrode surface. As a consequence, the potential energy profiles of reagents, intermediates, and products at the catalytic surface are lowered with respect to the noncatalytic electrode (Figure 24.5). This is mainly due to adsorption, especially in the case of ArBr • − and the reduction products Ar • and Br−, which strongly interact with the Ag surface, whereas for the starting reagents (ArBr + e−), their potential energy profile is lowered mainly because now the process occurs at less negative potentials. The combination of the decrease of the Gibbs free energy of the reaction and the enhancement of the cleavage rate of the radical anion results in a positive shift of E p at the catalytic electrode with respect to the noncatalytic one. In some cases, there is a change of mechanism from stepwise to concerted on passing from the noncatalytic electrode to the catalytic one (Figure 24.6). As before, the potential energy profiles of the reagents, intermediates, and products of the catalytic process are lowered by surface interactions and by the increase of the reduction potential at the catalytic electrode. It is worth noting that the change of DET mechanism from stepwise at GC to concerted at Ag causes a drastic change of both the standard free energy, ΔrG 0, and the activation free energy, ΔG ≠, of the ET. 0 0 It is well recognized that Estepwise is considerably more negative than Econcerted , even without taking into account adsorption phenomena, which would shift the latter to more positive values because of the preferential adsorption of the products with respect to the reagents. On the other hand, since the concerted process involves rupture of a chemical bond, its activation free energy is often greater than that of the stepwise process. Catalysis arises from a combination of these kinetic and thermodynamic effects, both related to the ET to ArBr. The process at Ag has a remarkable

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931

Activation of the Carbon–Halogen Bond OHP

IHP RBrads +e– – RBr ads

Ag electrode

+e–

Rads + Br–ads

+e–

Br–ads+ R–

R– + Br–

SCHEME 24.3 Reduction mechanism of aryl bromides at catalytic electrodes. (Reprinted from Isse, A.A. et al., J. Phys. Chem. C, 113, 14983, 2009. With permission.)

RX

Free energy

RX + e–

ΔrGθ

–

– RXads

ΔrGθ

R + X– RXads

+ e– Rads + X–ads Reaction coordinate

FIgURE 24.5 Potential energy profiles of reagents, intermediates, and products for aryl bromide reduction at the noncatalytic (full lines) and catalytic surface (dashed lines). (Reprinted from Isse, A.A. et al., J. Phys. Chem. C, 113, 14983, 2009. With permission.)

thermodynamic advantage over ET at a noncatalytic electrode but loses something from the kinetic standpoint. The overall result, however, favors the catalytic surface with a net positive shift of the reduction potential. A high electrocatalytic effect has also been observed for p-iodotoluene [128]. In this case, there is undoubtedly an important thermodynamic contribution due to the specific adsorption of I− at Ag. On the other hand, in this case even at the noncatalytic electrode, the DET mechanism is in a

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Organic Electrochemistry

RX

–

RX + e– –Fη Free energy

ΔrGo

– RX ads

–Fη

R + X– RXads + e–

ΔrGo

Rads + X–ads Reaction coordinate

FIgURE 24.6 Potential energy profiles of reagents, intermediates, and products for aryl bromide reduction at the noncatalytic (thin lines) and catalytic surface (thick lines), in the case of DET mechanism change from stepwise to concerted. (Reprinted from Isse, A.A. et al., J. Phys. Chem. C, 113, 14983, 2009. With permission.)

borderline situation [198] and it is highly likely to shift to the concerted limiting situation as the driving force of the reaction decreases owing to the electrocatalytic effect of Ag; this appears consistent with the peculiar peak sharpness observed at Ag, accounting for a determining adsorption process, which had been so far observed only for halides undergoing reduction according to the concerted mechanism. As discussed before, the nature of the halogen atom and the extent of charge localization on the C–X bond modulate the ability of the incipient leaving group to interact with the silver surface at a given potential, determining the extent of the catalytic effects. A further molecular feature that can enhance such interactions is the presence of adsorption auxiliary groups, having themselves specific affinity for the surface but undergoing no ET processes at the operating potential and therefore acting as anchoring groups, keeping the reacting C–X bond close to the surface. An example is the presence of a sulfur atom: Ag exhibits a higher catalytic activity for 2-bromothiophene than for bromobenzene, although both aryl halides reduce at similar potentials with a similar stepwise mechanism [136]. This finer molecular effect is somehow overshadowed by the high catalytic activity of silver but becomes much more relevant in the case of Au. This metal has for bromobenzene a lower and less reproducible catalytic effect than Ag on account of the repulsive effect of its very negative surface charge in the working potential range (see Section III.C.1). However, it approaches Ag activity in the case of 2-bromothiophene, where the anchoring S group is adjacent to the Br group to be cleaved. The beneficial anchoring effect is lower when it has to be shared between two Br leaving groups adjacent to the S group and becomes negligible in the case of a bromide leaving group in the 3-position [136]. 4. Electrocatalytic Effects of Silver: Role of Surface Morphology The surface morphology modulates the electrocatalytic effects of silver for organic halide reduction, although to a moderate extent. An exhaustive comparative investigation has been carried out on the reduction of several model halides on Ag(111), (110), and (100) monocrystals and on controlled polycrystalline silver surfaces of increasing roughness, in acetonitrile + 0.1 M tetraethylammonium perchlorate [166]. The reduction potentials of the alkyl, glycosyl, and benzyl halides are significantly shifted in the positive direction with increasing surface roughness (for polycrystals) or atomic

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Activation of the Carbon–Halogen Bond

933

densities and/or surface faceting (for monocrystals, indicating the following trend for the catalytic properties of the metal: 110 < 100 < 111). The Ep span for a given molecule on the different surfaces tested appears to be larger for molecular structures with higher electrocatalytic effects on silver; in particular, the maximum, 0.12 V, was found for acetobromoglucose. Evaluation of the effective surface area of the aforementioned polycrystalline Ag electrodes with respect to GC is important for the elucidation of the reaction mechanism in the catalytic case. To achieve it, overcoming the problem of the mechanism change on the catalytic surface, a procedure has been proposed [199], based on the comparative analysis of the CV features of a suitable set of probe molecules on both the catalytic and noncatalytic electrodes and evidencing inter alia the different surface perspectives of molecules reacting at the IHP rather than at the OHP. Of course, all the aforementioned observations hold for clean surfaces; filming by sparingly soluble reaction products or specifically adsorbed ions from the supporting electrolyte can result in signal irreproducibility and in significantly hindering the surface catalytic properties. In situ surface regeneration can be sometimes achieved by appropriate potential sweeps [106]. In view of industrial applications, some silver alloys were investigated in comparison to bulk silver to improve lifetime and shelf life of the electrode material [115]; all the tested alloys were found to be less electrocatalytic than silver, but some AgBi and AgSn alloys can be considered promising alternatives. The electrocatalytic activity of microsized silver powders supported on cavity microelectrodes was also investigated [200]. These supports are quite attractive because of their low impact on the supported materials (they require neither special manipulations nor sticking agents) and of the offered possibility of quick and reliable renovation of the electrode surface. These innovative Ag electrodes gave better performances than Ag electrodes prepared according to the conventional electrodeposition procedures, in terms of improvement of the electrocatalytic activity, insignificance of ohmic drop and double-layer capacitance in the voltammetric response, and the simplicity offered by the experimental procedure for renovating the electrode material and surface. More recently, several research groups tested the electrocatalytic activity of Ag nanoparticles [86,137]. Stable Ag nanoclusters were deposited on GC by a single potential pulse method in CH3CN + 0.1 M LiClO4 containing millimolar amounts of AgClO4. The particles, obtained by applying a pulse from rest potential to −0.4 V versus Ag|Ag+, are spherical in shape and are uniformly distributed over the GC surface with areal density number and particle size depending on applied deposition potential and deposition time, respectively. These Ag nanoclusters show comparable activity to that of the bulk metal for a wide range of compounds. Ag/GC electrodes of suitable active surface can be employed in macroscale electrolyses for important catalytic electrosyntheses. For example, electrocatalytic reduction of benzyl chloride gives satisfactory yields of toluene or phenylacetic acid, the latter being the main product obtained in electrocarboxylation conditions, that is, in CO2-saturated CH3CN. It is important to note that the performance of the Ag-modified GC electrode described here is comparable to that of bulk silver, but the catalyst metal load is very different, only a few micrograms of Ag per cm2 being present in the former. This is a very important economic aspect, which should be taken into consideration in large-scale electrosynthesis processes. In the same perspective, a study has been carried out [118] on silver nanoparticles (Ag_NP), synthesized by chemical reduction of an aqueous silver salt in the presence of six different stabilizing agents and supported on carbon powder (10% loading) for further characterization and use, and gasdiffusion electrodes (GDEs) based on the most promising Ag_NP composite have been successfully tested in an electrolytic process for the progressive conversion of gaseous trichloromethane to less chlorinated compounds and ultimately to methane. 5. Electrocatalytic Effects of Silver: The Role of the Supporting Electrolyte a. Nonspecifically Adsorbed Anions When both the supporting electrolyte ions have no specific affinity for the catalytic surface, double-layer effects should be determined by the supporting electrolyte cation, usually a quaternary

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Organic Electrochemistry

ammonium ion when operating in nonaqueous solvents, which must be located on the OHP according to the negative operating potentials with respect to the pzc. A comparative rationalization has very recently been proposed on the effect of quaternary ammonium supporting electrolyte cations (with C2, C3, C4, C6, and C8 alkyl chains), at constant counter anion, in aryl halide reduction on GC and silver electrodes [171]. On both electrode surfaces, increasing the electrolyte cation length results in a regularly increasing hindering effect, evident from both the negative potential shift and the widening of the reduction peaks. This effect, which is significantly higher on Ag than on GC in spite of the less extreme operating potentials, is remarkably dependent on the aryl halide molecular structure, regularly increasing with decreasing κ, that is, with the ET barrier becoming increasingly more determining with respect to the chemical one. In the case of the noncatalytic GC, the effect can be rationalized in terms of increasing doublelayer thickness, resulting in an increased hindrance of the electron tunneling between electrode and molecule, an effect becoming more and more determining with decreasing κ (which appears to be related to the β parameter in the electron tunneling probability equation), whereas in the catalytic Ag case, it has been proposed to depend on the increasingly smaller effective potential difference available to the molecule reacting close to the electrode when the double-layer thickness increases as a consequence of increasing the bulkiness of TAA+ cations [171]. b. Specifically Adsorbed Anions If supporting electrolytes containing halide anions are used, the latter can be specifically adsorbed on silver in a significant portion of the usual operating potential range for organic halide reductions, as discussed previously (see Section III.C.2) [127,166,168,195]. This results in a screening effect evidenced by a significant shift of the peak potential in the negative direction with respect to a given nonspecifically interacting medium, corresponding to an additional term in the activation energy accounting for halide anion desorption and increasing with both the halide anion affinity for silver (Cl− < Br− < I−) and the halide concentration. In experiments done with increasing I− concentrations, strikingly linear Ep versus log cI− characteristics are obtained, both on poly- and monocrystalline surfaces, at least in the case of the strongly adsorbed iodide. Such linearity and the relevant slopes (in terms of both negative sign and quasi-Nernstian values) are consistent with classic equations (Frumkin, Esin–Markov) describing simple electrochemical adsorption/desorption equilibria under the condition of constant surface coverage or charge. This indicates that the halide desorption step proves determining on the overall process, so the faradaic Ep values at each cI− can be regarded as indicators of a particular surface state, allowing the faradaic process to take place. In the case of the less strongly adsorbed chlorides and bromides, the same considerations apply, not surprisingly, only to a restricted, relatively high concentration range, below which such halide anions prove nearly ineffective. Another key feature of the Ep versus cX− characteristics is that all of them tend to an asymptote for concentrations as high as 0.1 M, which is perfectly reasonable in dealing with a surface phenomenon, that is, implying that a saturation condition must be reached. Accordingly, the maximum shift of ΔEp increases in the Cl− < Br− < I− sequence (50, 130, and 250 mV, respectively). They can be correlated to the logarithms of their respective equilibrium constants for chemical adsorption onto silver and exhibit a strikingly linear correlation with the logarithms of the solubility products of the corresponding silver halides. This implies, reasonably, that the halide adsorption constants on silver are proportional to the solubility constants of the corresponding silver halide. 6. Electrocatalytic Effects of Silver: The Role of the Solvent A recent systematic study shows that the catalytic effects of Ag for organic halide reduction hold, with the earlier described general features, both in aprotic (MeCN, DMF, PC, ACE, DMSO) and in protic solvents (BuOH, PrOH, EtOH, MeOH, H2O) [120,135,159,170]. However, while in the first group the catalytic effects appear nearly constant, in the second one they are higher and regularly

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Activation of the Carbon–Halogen Bond

increase with the solvent proticity; moreover, they appear to linearly increase with increasing primary medium effect (PME) on the halide leaving group, that is, the transfer activity coefficient of the halide leaving group from water to a solvent S [201] log γ t(X− ,W →S), a quantity directly proportional to the standard Gibbs energy of transfer of the same ion from infinite dilution in water to infinite dilution in S: log γ t(X− ,W →S) =

∆Gt0( X− ,W →S)

(24.33)

(2.303RT )

The more positive the PME, the weaker the solvent coordinating ability for the halide anion compared to water. In the alcohol series, PMEs are lower than in aprotic solvents and regularly decrease with increasing proticity. Actually, Ag catalytic effects also exhibit a fairly linear dependency on the alcohol pKa (Ka = autoprotolysis constant). Solvent proticity not only enhances the catalytic effects (maybe on account of faster turnover of halide ions in the catalytic sites of the electrode), but it can also affect the reaction mechanism; for example, in protic solvents, silver shows some catalytic effects also in the case of aryl chloride reduction, whereas in aprotic solvents, aromatic C–Cl bonds are not catalytically reduced on Ag [86,128,129]. Moreover, adsorption/desorption processes of straight-chain alcohols, regularly modulated by their chain length, are perceivable when working in such media. Water (PME = 0) appears to be the limiting case, resulting in the highest catalytic effects observed for each model molecule tested in the aforementioned solvent series (often ≥1 V). This feature is particularly valuable for synthetic and analytical applications in the environmental field, typically involving the aqueous medium. In fact, working on noncatalytic electrodes in water, the background reaction overshadows most organic halide reduction peaks, while working on Ag, a large number of halide molecules can be reduced well before the background cathodic limit; moreover, the corresponding reduction peaks are much more differentiated since catalytic effects depend on molecular structures (see Section III.C.3). Accordingly, working on silver widely increases the range of organohalide pollutants that can be abated by direct electroreductive dehalogenation and that can be monitored by voltammetric sensors [120]. Extension of the present studies to ionic liquid media, acting as both solvent and supporting electrolyte, could open further interesting perspectives.

D. PALLADIUM, COPPER, GOLD Palladium and copper also exhibit a remarkable catalytic effect for carbon–halogen bond cleavage; as an example, they both appear comparable to silver in a systematic parallel test carried out on a series of organic chlorides, including chloroaromatics, benzyl chlorides, activated chloroalkanes, and polychloromethane (PCM) [86]. Palladium has been, so far, more extensively investigated. At smooth palladium electrodes, palladiated surfaces (such as Cu, Ni, Au, Pt, GC) [138,139,145,151,155,162,202], and above all palladized silver [150,203], a large palette of substrates, including alkyl iodides [138,139], alkyl bromides [173,174], aryl halides [204], benzyl halides [148], allyl bromides [153], propargyl bromides [205], and vinyl bromides [206], have been investigated in different organic solvents (DMF, propylene carbonate, MeCN) with tetraalkylammonium quaternary salts as supporting electrolytes. In all cases, two subsequent reduction CV peaks were observed, the first of which is shifted to remarkably more positive potentials with respect to noncatalytic Pt and GC electrodes. In preparative experiments, the monoelectronic stoichiometry and the radical coupling and cross-coupling products pointed to a radical rather than an anion intermediate (as in the silver case), which was also supported by ESR experiments [139,148]. An organometallic intermediate was assumed for the catalytic process, possibly leading to significant chemical modifications of the catalytic metal surface and even to its partial dissolution [207].

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936

Organic Electrochemistry

From a preparative point of view, the catalytic process resulted inter alia in arene alkylations [145], mono- and di-benzylations [148], formation of homo dimers from long-chain α,ω-dibromoalkanes [146], and aryl dimerizations in propylene carbonate (for which a probable concomitance of electrocatalytic and redox catalytic processes was assumed [208]). More recently, research also focused on copper [86,121–124,132,158] as smooth metal or in other forms (such as copper powder suspensions, palladium copper alloy, or galvanostatically electrodeposited copper onto several conducting substrates). Of applicative interest for environmental remediation processes is the copper performance in the electrocatalytic reductive dehalogenation of chlorinated compounds such as PCMs [121], geminal polychloroethanes (PCAs, in particular 1,1,1-trichloroethane (TCA) and 1,1-dichloroethane (DCA), which are the simplest molecules belonging to the homologous series of CHCl3 and CH2Cl2 [122]), and polychloroethylenes [124]. Although in DMF the Cu catalytic effects for the process are modest, they are significantly enhanced by the addition of proton donors, which also affect the nature of the intermediates and products in preparative experiments. In particular, sequential hydrodehalogenation leading to ethane as the final product becomes the principal reaction pathway in the presence of acetic acid, whereas in the presence of H2O both hydrodehalogenation and dehydrodehalogenation (with α,β-elimination of H+ and Cl− resulting in chlorinated olefins and acetylene) are possible. Gold is also currently studied [124,136,159,160–162], focusing on conditions in which its intrinsic catalytic properties can be exploited in spite of the high negative surface charge, such as in the presence of adsorption auxiliaries [136] or in protic solvents [159]. The electrocatalytic reduction of organohalides on the aforementioned metals has been proposed for electrode surface functionalization (e.g., [160–162]).

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Isse, A.A.; De Giusti, A.; Gennaro, A. Tetrahedron Lett. 2006, 47, 7735–7739. Isse, A.A.; Mussini, P.R.; Gennaro, A. J. Phys. Chem. C 2009, 113, 14983–14992. Savéant, J.-M. Acc. Chem. Res. 1980, 13, 323–329. Matyjaszewski, K.; Xia, J. Chem. Rev. 2001, 101, 2921–2990. Isse, A.A.; Gennaro, A. Chem. Commun. 2002, 2798–2799. Costa, G. Coord. Chem. Rev. 1972, 8, 63–75. Isse, A.A.; Gennaro, A.; Vianello, E. J. Chem. Soc., Dalton Trans. 1993, 2091–2096. Isse, A.A.; Gennaro, A.; Vianello, E. J. Chem Soc., Dalton Trans. 1996, 1613–1618. Isse, A.A.; Gennaro, A.; Vianello, E. J. Electroanal. Chem. 1998, 444, 241–245. Cardinale, A.; Gennaro, A.; Isse, A.A.; Maran, F. Substitution and dissociative electron transfer to benzyl halides. In New Directions in Organic Electrochemistry, Fry A.J., Matsumura Y., Eds.; The Electrochemical Society Proceedings Volume Series: Pennington, NJ, 2000, pp. 136–140. Isse, A.A.; Gennaro, A.; Vianello, E. Electrochim. Acta 1992, 37, 113–118. Gennaro, A.; Isse, A.A.; Maran, F. J. Electroanal. Chem. 2001, 507, 124–134. Isse, A.A.; Ferlin, M.G.; Gennaro, A. J. Electroanal. Chem. 2003, 541, 93–101. Esteven, A.P.; Goken, D.M.; Klein, L.J.; Lemos, M.A.; Medeiros, M.J.; Peters, D.G. J. Org. Chem. 2003, 68, 1024–1029. Tsuji, J. New J. Chem. 2000, 24, 127–135. Mizoroki, T.; Mori, K.; Ozaki, A. Bull. Soc. Chim. Jpn. 1971, 44, 581–581. Heck, R.F.; Nolley, J.P., Jr. J. Org. Chem. 1972, 37, 2320–2322. Cassar, L. J. Organomet. Chem. 1975, 93, 253–257. Dieck, H.A.; Heck, R.F. J. Organomet. Chem. 1975, 93, 259–263. Sonogashira, K.; Tohda, Y.; Hagihara, N. Tetrahedron Lett. 1975, 16, 4467–4470. Fauvarque, J.-F.; Jutand, A. Bull. Soc. Chim. Fr. 1976, 765–770. Negishi, E.-I; King, A.O.; Okukado, N. J. Org. Chem. 1977, 42, 1821–1823. Fauvarque, J.-F.; Jutand, A. J. Organomet. Chem. 1977, 132, C17–C19. Milstein, D.; Stille, J.K. J. Am. Chem. Soc. 1979, 101, 4992–4998. Miyaura, N.; Suzuki, A. J. Chem. Soc., Chem. Commun. 1979, 19, 866–867. Hatanaka, Y.; Hiyama, T. Tetrahedron Lett. 1988, 29, 97–98.

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Aliphatic and Aromatic Halides

25 Conversions Dennis G. Peters CONTENTS I. II.

Introduction .......................................................................................................................... 942 Monohalogenated Alkanes ................................................................................................... 942 A. Reduction at Mercury Cathodes ................................................................................... 943 1. Primary Alkyl Monohalides ................................................................................. 943 2. Secondary Alkyl Monohalides .............................................................................. 943 3. Tertiary Alkyl Monohalides .................................................................................. 943 B. Reduction at Carbon Cathodes .....................................................................................944 C. Reduction at Other Electrodes......................................................................................944 D. Electrolyte Effects ........................................................................................................ 945 III. Polyhalogenated Alkanes ..................................................................................................... 945 A. gem-Dihaloalkanes, gem-Trihaloalkanes, and Carbon Tetrachloride .......................... 945 B. Vicinal Dihaloalkanes and Trihaloalkanes ..................................................................946 C. α,ω-Dihaloalkanes ........................................................................................................ 947 IV. Halogenated Alkenes and Alkynes ...................................................................................... 947 A. Vinyl Halides ................................................................................................................ 947 B. Allyl Halides................................................................................................................. 947 C. Acetylenic Halides........................................................................................................948 V. Benzyl Halides and Related Compounds ............................................................................. 948 VI. Alicyclic Halides .................................................................................................................. 950 A. Monohalogenated Species ............................................................................................ 950 B. Dihalogenated Species ................................................................................................. 951 C. Other Species................................................................................................................ 951 VII. Halogenated Aromatic Compounds ..................................................................................... 952 A. Monohalobenzenes ....................................................................................................... 952 B. Polyhalobenzenes ......................................................................................................... 953 C. Halogenated Nitrobenzenes and Cyanobenzenes......................................................... 954 D. Halogenated Aryl Carbonyls ........................................................................................ 954 E. Halogenated Biphenyls ................................................................................................. 954 VIII. Acyl Halides ......................................................................................................................... 955 IX. Phenacyl Halides .................................................................................................................. 955 X. Aliphatic α-Halocarbonyl Compounds ................................................................................ 956 XI. Halogenated Heterocyclic Compounds ................................................................................ 958 A. Heterocyclic Compounds with One Nitrogen Atom .................................................... 958 B. Heterocyclic Compounds with Two Nitrogen Atoms ................................................... 959 C. Other Heterocyclic Species .......................................................................................... 959

941

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XII. Indirect or Catalytic Cleavage of Carbon–Halogen Bonds .................................................. 959 A. Electrogenerated Radical–Anions as Mediators ..........................................................960 B. Transition-Metal Complexes as Mediators ...................................................................960 1. Nickel Complexes .................................................................................................. 961 2. Cobalt Complexes .................................................................................................. 963 3. Other Catalysts ...................................................................................................... 965 C. Reduction of Carbon–Halogen Bonds at Chemically Modified Electrodes ................ 965 References ......................................................................................................................................966

I. INTRODUCTION Electrochemical reduction of halogenated organic compounds, including mechanistic features, as well as synthetic applications associated with reductive cleavage of carbon–halogen bonds, has been treated in earlier editions of this work [1–4] and in several other reviews [5–11]. This chapter focuses both on the classic material described in these earlier references and, more importantly, on new achievements published since 2000.

II.

MONOHALOgENATED ALKANES

Electrochemical reduction of a simple alkyl monohalide (RX) has been visualized classically in terms of the following mechanistic picture: RX + e

[RX –]

[RX –]

R + X– R–

R +e

Combination

R2

2R Disproportionation

RH + R(–H)

R

+

SH

RH

+

S

R–

+

SH

RH

+

S–

R–

+

RX

R2

+

X–

R–

+

HB

RH

+

B–

RX +

B–

RX +

B–

E2

SN2

R(–H) + X– + HB

RB

+

X–

However, recent publications by Isse and coworkers [12–14] have offered new perspective on whether the radical–anion of RX actually exists and on other details about the cleavage of carbon– halogen bonds (see Chapter 24). Nevertheless, depending on the solvent, supporting electrolyte,

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electrode material, and potential, it is possible to electrogenerate either alkyl radicals or carbanions that can lead to the formation of dimers (R2), alkanes (RH), and olefins [R(–H)]. In addition, the solvent (SH) can act as a hydrogen atom donor or a proton donor. Also, olefins can arise from the base-promoted dehydrohalogenation of RX, and other products (RB) can be formed via displacement of halide from RX by a base (B−). When cathode materials such as mercury, lead, and tin are employed, electrogenerated alkyl radicals can interact strongly with the electrode to enhance the reduction of the alkyl monohalide and to afford organometallic compounds. Tertiary alkyl halides are easier to reduce than secondary alkyl halides, which are, in turn, easier to reduce than primary alkyl halides. Ease of reduction of a carbon–halogen bond depends on the identity of the halogen atom: (1) iodides are easier to reduce than bromides, (2) chlorides are so difficult to reduce that they often appear to undergo no direct reduction, and (3) direct reduction of an alkyl monofluoride has not been observed. Finally, the existence of the radical–anion [RX• –], formed by the addition of one electron to RX, has not been demonstrated; Andrieux and coworkers [15] discussed why the radical–anion is not expected for simple alkyl monohalides, whereas radical–anions of aromatic halides are distinct intermediates in the electrochemical reduction of aryl halides.

A. REDUCTION AT MERCURY CATHODES 1. Primary Alkyl Monohalides Many studies of the reduction of primary alkyl monohalides at mercury electrodes have been undertaken. Reduction of iodoethane and 1-iodobutane in ethanol affords diethylmercury and dibutylmercury, respectively, in essentially quantitative yields [16]. When 1-iododecane is reduced at mercury in DMF, successive formation of the decyl radical (to afford only didecylmercury) and the decyl carbanion (to yield mainly decane and 1-decene) takes place [17,18]. Bilewicz and Osteryoung [19] found that the reductions of iodoethane, 1-iodobutane, and 1-iododecane at mercury in MeCN are similar to those seen in DMF. Primary alkyl monobromides, such as 1-bromodecane, typically undergo a one-step, twoelectron reduction to alkyl carbanions in DMF, a process leading to alkanes and alkenes [18]. However, Fry [20] reported that reduction of 1-bromohexane in DMF containing TEABr yields hexane, 1-hexene, dihexylmercury, and 1-hexanol. Wagenknecht and coworkers [21,22] investigated the reductions of bromoethane and 1-bromobutane at mercury in carbon dioxide–saturated DMF. For the 1-bromobutane–carbon dioxide system, dibutylmercury, dibutyl oxalate, and butyl valerate are produced in comparable yields; to account for the formation of the dialkylmercury compound, the intermediacy of a short-lived adsorbed alkylmercury radical is proposed. 2. Secondary Alkyl Monohalides One detailed report [23] about the reduction of secondary alkyl monohalides at mercury has appeared. In DMF containing tetraalkylammonium salts, 2-iodooctane undergoes stepwise reduction, whereas 2-bromooctane exhibits just one stage of reduction. One-electron reduction of 2-iodooctane affords the sec-octyl radical, which leads to a mixture of octane, 1-octene, 2-octene, 7,8-dimethyltetradecane, and di-sec-octylmercury. Products obtained from the two-electron reduction of either 2-bromo- or 2-iodooctane are as just mentioned, but the yield of octane is higher and the amounts of di-sec-octylmercury and 7,8-dimethyltetradecane are lower. Montero and coworkers [24] found that reduction of 2-bromo-2-nitropropane at mercury in dichloromethane containing dimethyl but-2-ynedioate gives two unexpected products—namely, dimethyl 2-(2-nitropropan-2-yl)but-2-enedioate and trimethyl 2,2-dimethyl-2H-pyrrole-3,4,5-tricarboxylate. 3. Tertiary Alkyl Monohalides Polarographic and voltammetric investigations [25–28] of the behavior of tert-butyl iodide and tertbutyl bromide at mercury revealed that, depending on the solvent and supporting electrolyte, the iodide undergoes stepwise reduction and the bromide exhibits either a single two-electron reduction

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or a pair of one-electron reductions. Bulk electrolyses of tert-butyl bromide at potentials corresponding to the first wave produce tert-butyl radicals, which form isobutane, isobutylene, and 2,2,3,3-tetramethylbutane; however, at more negative potentials, where tert-butyl carbanions are generated, one obtains isobutane and isobutylene in equal yields [29].

B. REDUCTION AT CARBON CATHODES An advantage of carbon electrodes over mercury cathodes for the reduction of alkyl monohalides is the avoidance of dialkylmercury species. As pointed out by Lambert and Ingall [30], it is possible with carbon cathodes to observe the direct reductive cleavage of the carbon–chlorine bond in compounds such as 1-chloropentane, 1-chlorohexane, 2-chlorobutane, and tert-butyl chloride in DMF. Reduction of 1-bromobutane at graphite in carbon dioxide–saturated DMF is a two-electron process that affords butane, 1-butene, octane, butyl valerate, butyl-N,N-dimethyloxamate, and dibutyl-2-methylmalonate [22]. Kaabak and coworkers [31] found that reduction of iodoethane and 1-bromobutane at graphite in DMF yields alkanes, olefins, dimers, and solvent-derived products. Using glassy carbon in DMF containing TBABF4, Andrieux and coworkers [15,32] concluded that 1-bromobutane, 2-bromobutane, tert-butyl bromide, and 1-iodobutane undergo one-step, twoelectron reductions, whereas 2-iodobutane and tert-butyl iodide exhibit a pair of one-electron reduction waves. Products derived from bulk electrolyses of primary, secondary, and tertiary alkyl iodides and bromides at reticulated vitreous carbon cathodes have been determined in DMF [33]. In work involving the use of reticulated vitreous carbon cathodes in DMF [34], the reductions of iodoethane and 2-iodopropane were examined and product distributions were reported. After it was discovered that reduction of secondary alkyl halides (RR′CHX) causes RR′CH– groups to be grafted onto carbon surfaces [35] (see also Chapter 42), Jouikov and Simonet [36] verified this process by reducing ferrocene bearing an alkyl iodide substituent to graft alkylferrocenyl moieties as probes of this surface phenomenon. Another paper deals with the grafting of pendant alkanoic acid moieties onto carbon via the reduction of ω-bromoalkanoic acids [37]. Gennaro and coworkers [38] report that changes in the oxygen-to-carbon ratio on the surface of glassy carbon plays an important role in the reduction of carbon–halogen bonds.

C. REDUCTION AT OTHER ELECTRODES Methyl, ethyl, propyl, and butyl halides are reduced at lead and tin cathodes to produce, respectively, the corresponding tetraalkyllead and tetraalkyltin compounds in excellent yield. Bismuth, gallium, indium, and thallium can be employed to prepare a variety of organometallic compounds. Feoktistov [2] and Hawley [5] have summarized many of these electrosyntheses. Sock et al. [39] synthesized carboxylic acids by reducing alkyl bromides at platinum, gold, stainless steel, and graphite in the presence of carbon dioxide in either DMF or a THF–HMPA mixture. Reduction of an alkyl monobromide at nickel in DMF and in the presence of an arylalkene and a sacrificial aluminum anode leads to an addition product [40]. Simonet and coworkers [41–43] discovered that, when platinum cathodes are employed for the direct reduction of alkyl halides in superdry DMF containing tetramethylammonium tetrafluoroborate, the original alkyl halide is converted into an alkene and the surface of the platinum electrode is transformed into an ionometallic layer. A host of reports have appeared that focus on the reduction of alkyl monohalides at nontraditional electrodes such as palladized carbon, copper, nickel, or platinum; copper–palladium or silver–palladium alloys; and silver [44–55]. The emphasis in these papers is on the ability of the electrode to promote the facile one-electron scission of the carbon– halogen bond. Isse et al. [56] have compared mechanisms for the reduction of alkyl and aromatic halides at glassy carbon and silver, pointing out that silver exhibits an amazing ability to catalyze cleavage of the carbon–halogen bond; similar effects have been seen with palladium and copper electrodes [57]

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(see also Chapter 24). Dissociative electron-transfer processes, as they influence the reductive cleavage of carbon–halogen bonds at silver cathodes, have been studied by the same group of workers [58]. Synthesis of cyanoacetic acid via reduction of haloacetonitriles in the presence of carbon dioxide has been compared for silver and glassy carbon electrodes, and important mechanistic conclusions have emerged [59–62]. A comparison of the reduction of various bromonitriles at both mercury and glassy carbon was undertaken [63]. Reactions of Br(CH2)nCN (n = 2–4) with the anion arising from the reduction of 2-bromo-2-methylpropanenitrile have been investigated [64], and the reduction of a series of bromonitriles at mercury and glassy carbon has been studied [65]. Monoand polycrystalline silver electrodes have been employed to study the reductions of alkyl halides (as well as benzyl and glycosyl halides) [63,64].

D.

ELECTROLYTE EFFECTS

For electrosyntheses, it is important to be aware of the substantial shifts in the potentials needed to reduce carbon–halogen bonds when different supporting electrolytes are used. Many workers [18,19,23,27,28,68] have noticed that carbon–halogen bonds become more difficult to cleave electrolytically as the size of the tetraalkylammonium cation of the supporting electrolyte increases. Fry and Krieger [27] determined the relative electron-transfer rates in DMSO for the reduction of some alkyl and aryl halides as a function of the identity of the tetraalkylammonium cation comprising the supporting electrolyte. Alkyl carbanions are potent bases, so they are protonated by almost any source of hydrogen ions, especially adventitious water in the solvent or supporting electrolyte. In the absence of water and other added proton donors, tetraalkylammonium cations can serve as proton donors toward alkyl carbanions; deprotonation of tetraalkylammonium ions leads, via the Hofmann elimination, to the corresponding trialkylamine and olefin. When tetramethylammonium salts are used as supporting electrolytes, there is evidence that trimethylammonium methylide is formed [69,70].

III.

POLyHALOgENATED ALKANES

A.

gem-DIHALOALKANES, gem-TRIHALOALKANES, AND CARBON TETRACHLORIDE

Electroreduction of polyhalogenated alkanes is a useful strategy for organic synthesis and environmental remediation. Among studies in this area are the following: reduction of polychloromethanes and polychloroethanes at activated silver cathodes [71–73]; selective dechlorination of 1,1-dichloro-2,2-bis(p-chlorophenyl)ethane and 1,1-dichloro-2-hydroxy-2,2-bis(p-chlorophenyl)ethane [74]; formation of cyclopropanes via reduction of dichloromethane in the presence of alkenes [75,76]; cross-coupling of activated alkyl halides with gem-trichloroalkanes or carbon tetrachloride [77,78]; electrochemical coupling of gem-trichloroalkanes with aldehydes or ketones [79]; preparation of cyclopropanes by indirect electroreductive coupling of activated olefins and gem-dichloro or gem-trichloro compounds [80]; synthesis of epoxides via indirect reductive coupling of carbonyl compounds with activated gem-dichloro species [81]; conversion of tetra- and trichloropentanes into less chlorinated analogues [82]; reduction of 1,1,1-trichloro-2-hydroxy-4methyl-4-pentene to afford 1,1-dichloro-4-methyl-2,4-pentadiene and 1,1-dichloro-2-hydroxy-4methyl-4-pentene [83]; electrolysis of trichloromethylphosphonate in the presence of an alkyl halide to form a diethyl 1,1-dichloroalkylphosphonate [84]; reductive coupling of 1,1,1-trichloro-2,2,2trifluoroethane, methyl chlorodifluoroacetate, trifluorobromomethane, and perfluoroiodobutane with aldehydes [85]; reductive cyclization of 2,2-dimethyl-3-oxopropyl 2,2,2-trichloroacetate to afford 3,3-dichloro-4-hydroxy-5,5- dimethyltetrahydro-2H-pyran-2-one [86]; synthesis of 3-chloro-1-methyl-4-phenylquinolin-2(1H)-one from N-(2-benzoylphenyl)-2,2,2-trichloroN-methylacetamide [87]; reductive cyclization of 2-acetylphenyl 2,2,2-trichloroacetate to produce 3,3-dichloro-4-hydroxy-4-methylchroman-2-one [88]; reductive dechlorination of

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4,4,4-trichloro-1-phenylbut-2-en-1-ones [89–91], N-(1-alkoxy-2,2,2-trichloroethyl)benzamides [92,93], N-(1,2,2,2-tetrachloroethyl)benzamides [94], 2,2,2-trichloroethylideneacetophenones [95], and 3,3,6,6-tetrachloro-1,2-cyclohexanedione [96,97]; dechlorination of alkyl and aryl carbonimidoyl dichlorides to give the corresponding isocyanides [98]; electroreductive coupling of activated olefins and gem-polyhalo compounds [99]; selective removal of gem-polyhaloethoxy protecting groups from carboxylic acids [100]; and electrochemical conversion of 2,2-dichloroethylamides to 4,5-dihydrooxazoles [101,102]. An example of the last conversion is the reduction of N-(2,2-dichloro1-phenylethyl)benzamide to give 2,4-diphenyl-4,5-dihydrooxazole in 78% yield: O

Ph

Ph

O

2e, –2Cl–

Ph

N H

CHCl2

Hg cathode CH3CN, TBAClO4

N Ph

Carbon tetrachloride undergoes stepwise reduction at mercury in DMF [103]. Several groups of workers [104–110] have electrogenerated the trichloromethyl anion, which reacts with acrylonitrile, ethyl acrylate, diethyl fumarate, alkyl monohalides, aldehydes, and ketones. Dichlorocarbene, formed via reduction of carbon tetrachloride, reacts with 2,3-dimethylindole to afford 3-chloro2,4-dimethylquinoline and 3-(dichloromethyl)-2,3-dimethyl-3H-indole [111]; electrolysis of carbon tetrachloride or tetrabromide in the presence of aryl aldehydes gives benzal chlorides and bromides, respectively [112]. Reduction of bromotrifluoromethane at stainless steel in cells with sacrificial anodes of cadmium, copper, and zinc can be used to prepare CF3Cd, CF3Cu, and CF3Zn [113–115]. Reduction of carbonimidoyl dichlorides at mercury gives isocyanides in excellent yield [116]. Reduction of chlorodifluoromethyl enol ethers results in two-electron cleavage of the carbon–chlorine bond, and the resulting anion can react with trimethylchlorosilane to afford functionalized difluoromethyl allylsilanes [117]. Comparative studies of the dehalogenation of polychlorinated methanes at glassy carbon and silver cathodes have been recently reported [118,119].

B.

VICINAL DIHALOALKANES AND TRIHALOALKANES

These compounds undergo a two-electron reduction with the loss of two halide ions to give an olefin. Casanova and Rogers [120] demonstrated the reductions of 1,2-dibromo-2-phenylpropane to  2-phenylpropane, of meso-2,3-dibromobutane to trans-2-butene, and of d,l-2,3-dibromobutane to cis-2-butene. For the reduction of d,l-1,2-dibromo-1,2-diphenylethane at mercury in DMF, the ratio of cis- to trans-stilbene seems to depend on the identity and concentration of the supporting electrolyte [121,122]. Reduction of vicinal dibromides has been carried out by Závada and coworkers [123] and by Lund and coworkers [124]. Papers concerning the reduction of polyhalogenated ethanes have been published by Feoktistov, Gol’din, and coworkers [125–129]. Rampazzo and coworkers [130] studied the reductions of α,α,α′,α′-tetrabromo-o-xylene and 1,2-dibromobenzocyclobutene. Various mechanistic aspects of the electroreduction of vicinal dihalides have been probed [120,131–135]. In an investigation of the meso- and d,l forms of 3,4-dibromohexane and 2,5-dimethyl-3,4-dibromohexane, Brown and coworkers [136] found that the meso compounds are easier to reduce than the d,l species. Evans and coworkers [137,138] employed cyclic voltammetry to probe the temperature-dependent conformational equilibrium for a number of vicinal dibromides, and Lexa and coworkers [139] have discussed inner- and outer-sphere processes for the reduction of vicinal dibromides.

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Aliphatic and Aromatic Halides

C.

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α,ω-DIHALOALKANES

In early work, Rifi [140–142] investigated the reduction of α,ω-dibromoalkanes (such as 1,3-dibromopropane and 1,4-dibromobutane) at mercury in DMF. Studies of α,ω-dibromoalkanes were conducted by Wiberg and Epling [143] and by Fry and Britton [144,145], and Casanova and Rogers [146] synthesized dialkylmercury species via controlled-potential electrolysis of a family of compounds (ranging from 1,4-dibromobutane to 1,12-dibromododecane) at mercury in DMF. A number of other publications have dealt with the behavior of α,ω-dibromoalkanes at mercury cathodes [147–152]. Derivatives of cyclopentane and cyclohexane have been prepared by the reduction of 1,3-dibromopropane or 1,4-dibromobutane, respectively, in the presence of an alkene [40,75]. In a series of papers pertaining to the reduction of α,ω-dihaloalkanes at carbon in DMF, the behavior of 1,3-dihalopropanes [153], 1,4-dihalobutanes [154], 1,5-dihalopentanes [155], 1,6-dihalohexanes [156], 1,8-dihalooctanes [157], and 1,10-dihalodecanes [157] was probed. Simonet [158] has investigated the reduction of several α,ω-dihaloalkanes at copper cathodes. Tokuda and coworkers [159] carried out reductive intramolecular cyclization of compounds such as 1,12-dibromo-2,10-dodecadiene to form the corresponding 1,5-cycloalkadienes.

IV. HALOgENATED ALKENES AND ALKyNES A.

VINYL HALIDES

Reduction of a vinyl halide involves the uptake of one electron and the loss of a halide ion to give a vinyl radical, which then undergoes further reduction and protonation to yield an alkene; before accepting an electron and a proton, the vinyl radical rearranges to afford either a cis or trans alkene. Compounds that have been specifically investigated include aryl-substituted vinyl bromides [160], cis- and trans-3-iodo-3-hexene [161], bromomaleate and bromofumarate [162], and aryl-substituted 3-chloroacrylonitriles [163]. Reductive carboxylation of phenyl-substituted vinyl bromides affords α,β-unsaturated carboxylic acids [164]. Yoshida and coworkers [165] electrolyzed vinyl halides at platinum in DMF in the presence of trimethylchlorosilane to obtain silylation products. Vicinal dibromoalkenes and diiodoalkenes undergo reduction, with the loss of both halide ions, to yield alkynes that are further reduced to olefins and saturated species; work by Rosenthal and coworkers [166] on the electrochemistry of diethyl dibromofumarate and diethyl dibromomaleate exemplifies this behavior. Vicinal dichloroalkenes show more complicated electrochemistry; in some instances, the two carbon–chlorine bonds simply undergo successive two-electron, one-proton reductive cleavage [167], whereas in other cases the reduction leads to an alkyne [168]. Feroci and coworkers [169] converted substituted 1,1-dibromoalkenes, for example, 2-(2,2-dibromovinyl)naphthalene, to vinyl bromides via electrolysis at a carbon, gold, mercury, or silver cathode; the E/Z ratio for the products depends on the choice of electrode.

B. ALLYL HALIDES Baizer and coworkers [104,170] demonstrated that reduction of allyl bromide and allyl chloride depends on the medium. In DMSO containing TEAP, allyl bromide undergoes stepwise reduction to the radical and to the carbanion, whereas allyl chloride exhibits a single two-electron reduction [170]. In DMF containing TEAOTs or TEABr, allyl bromide and allyl chloride are each reduced in a one-step, two-electron process to the carbanion, which then undergoes various follow-up reactions with a proton donor, with unreduced starting material, with carbon dioxide, or with species such as acrylonitrile and ethyl acrylate [104]. Reduction of allyl halides, 1-chloro-3-methyl-2-butene, and methyl-4-halo-2-butenoates at platinum in DMF containing diethyl fumarate gives the conjugate addition products [171]; and reduction of 4-bromo- and 4-chloro-2-butene can be used to allylate acetone and benzaldehyde [172]. Methallyl chloride undergoes reductive coupling with ketones and

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aldehydes to afford alcohols [173] as well as electrodimerization [174]. Reduction of allylic halides in the presence of anhydrides leads to the formation of ketones [175], as seen in the following process: H2C

2e, –Cl– Ni cathode DMF, TBAI

CH2Cl

C

C

H2C

CH3

CH3

O H2C

C CH3

CH2

CH2–

O

–

H3C

O

CH3

H2C

–CH3COO–

C CH3

CH2

C

CH3

O

Bard and Merz [176] reported that allyl bromide and allyl iodide interact chemically with a mercury electrode to form allylmercury halides, which are reduced to electroactive diallylmercury. Allyl bromide and allyl iodide are reduced at platinum in MeCN in a two-electron process to give the allyl anion. Reduction of allyl halides at platinum in DMF containing trimethylchlorosilane produces silylated compounds [165]. Simonet [177] has shown that reduction of allyl bromide at noble metals (Au, Pd, Pt, and Rh) results in the surface allylation of those electrodes. Electrochemical conjugate addition of the allyl moiety of a substituted allyl halide to an α,β-unsaturated ester has been described [171,178]. Electrolysis of allyl chloride in the presence of excess acetone gives 2-methyl-4-penten-2-ol [179]. Reduction of an allyl halide in the presence of copper(II) acetylacetonate leads to the formation of biallyl [180]. According to Brillas and Costa [181], 1,3-butadiene is an intermediate in the reduction of trans1,4-dibromo-2-butene and trans-1,4-dichloro-2-butene at mercury. Doupeux and Simonet [182,183] studied the electrochemical behavior of polyhalogenated allyl halides.

C. ACETYLENIC HALIDES Several investigations [184–187] have dealt with acetylenic halides that undergo reductive intramolecular cyclization. Electrolysis of 6-bromo- or 6-iodo-1-phenyl-1-hexyne at mercury in DMF gives a potential-dependent product distribution [184]. Reduction of 1-iodo-5-decyne at mercury displays two waves, indicating stepwise formation of radical and carbanion intermediates, whereas reduction of 1-bromo-5-decyne gives a single wave, which corresponds to a net two-electron process [185]; bulk electrolyses of these compounds afford potential-dependent product distributions. When 1-halo-5-decynes are electrolyzed at carbon in DMF [186], the yield of pentylidenecyclopentane increases dramatically; two other products are 5-decyne and 1-decen-5-yne. Reduction of 6-iodo-1-phenyl-1-hexyne at carbon in DMF affords benzylidenecyclopentane and 1-phenyl-1hexyne [187].

V. bENZyL HALIDES AND RELATED COMPOUNDS Reduction of benzyl iodide at mercury appears to proceed via an electroinitiated chain reaction that involves the formation of benzylmercuric iodide, C6H5CH2HgI [188,189]; benzyl bromide and benzyl chloride probably exhibit similar behavior [190–192]. Electrolyses of benzyl bromide in MeCN [192] or DMF [104] and of benzyl chloride in aqueous EtOH [192] at mercury afford dibenzylmercury, toluene, and sometimes bibenzyl; in the literature, there is much debate about mechanistic details of the reduction of benzyl halides [104,193]. Reduction of benzyl chloride at mercury in DMF and in the presence of carbon dioxide provides

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evidence for the generation of the benzyl carbanion; benzyl phenylacetate [104] and phenylacetic acid [189] are obtained as electrolysis products. In carbon dioxide–saturated MeCN, benzyl chlorides undergo facile electrocarboxylation at a silver cathode [194,195]; similar studies of the reductive carboxylation of benzyl bromide at graphite and mercury have been reported [196]. Benzyl chloride undergoes electrodimerization at a stainless-steel cathode in DMF [174]. Reduction of benzyl iodide at glassy carbon in the presence of nitric oxide has been investigated [197]. In the presence of copper(II) acetylacetonate, benzyl bromide undergoes reductive coupling to give bibenzyl [180]. Electrochemical reduction of several derivatives of 4-(bromomethyl)-2H-chromen-2-one at vitreous carbon has been investigated [198,199]. Electrolyses of 1-(2-chloroethyl)-2-nitrobenzene and 1-(2-bromoethyl)-2-nitrobenzene at glassy carbon electrodes in the absence of an added proton donor afford mainly 1-nitro-2-vinylbenzene; however, in the presence of an excess of a proton donor (e.g., phenol or 2,4-pentanedione), 1H-indole is obtained in good yield (>80%) [200]: X 4e NO2

C cathode, DMF, excess proton donor

N H

Reduction of N′-(2-bromo-1-phenylethylidene)acetohydrazide to (1E,4E)-1,4-bis(2-methylhydrazono)1,4-diphenylbutane has been reported [201]. Isse et al. [202–205] have compared the reduction of benzyl halides at silver and glassy carbon, and Amatore and coworkers [206,207] used electrochemistry along with spectroscopic and computational methods to elucidate the mechanism of the reduction of benzyl chloride at silver. Reduction of o- and p-nitrobenzyl halides at mercury or platinum gives mainly the bibenzyl derivative, with nitrotoluene as a minor product [193,208–212]. However, m-nitrotoluene is the predominant product of the electrolysis of m-nitrobenzyl halides [209]. Andrieux and coworkers [213] probed the mechanism of reductive cleavage of a family of substituted benzyl halides. Reduction of 4-nitrobenzyl bromide and benzyl bromide at highly ordered pyrolytic graphite results in the creation of a film of benzylic moieties on the cathode surface [214]. Koch and coworkers [215] investigated the reduction of benzyl halides at both platinum and carbon in MeCN. For the electrolysis of benzyl iodide at platinum, the products are toluene, bibenzyl, and hydrocinnamonitrile; in the presence of carbon dioxide or diethyl malonate, the electrogenerated benzyl anion is converted, respectively, into benzyl phenylacetate or the diethyl ester of benzyl malonate. Shono and coworkers [216] electrolyzed benzyl chloride or (1-chloroethyl)benzene at carbon in MeCN containing various acyl chlorides to prepare alkyl benzyl ketones. Ketones have been obtained from the reduction of benzyl halides in the presence of anhydrides [175]. Catalytic electrocarboxylation of (1-chloroethyl)benzene has been carried out [217], and 2-arylpropanoic acids have been synthesized via reduction of several 1-aryl-1-chloroethanes at silver and carbon cathodes in the presence of carbon dioxide [218]. Electrolysis of benzyl chloride in the presence of various ketones and aldehydes can be used to synthesize alcohols [110]: 2e, –Cl– PhCH2Cl

Stainless-steel cathode DMF, TBAI

PhCH2–

O

PhCH2–

H3C

C +H+

CH3 CH3

PhCH2

C OH

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CH3

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Organic Electrochemistry

For the reduction of 1-bromo-1-phenylethane at mercury in DMF, Yamasaki and coworkers [219] found a potential-dependent product distribution. At less negative potentials, the products are meso2,3-diphenylbutane, d,l-2,3-diphenylbutane, and phenylethane, whereas phenylethane is the only species obtained at more negative potentials. Using (1-bromo-2,2-dimethylpropyl)benzene, Fry and Powers [220] obtained more insight into the mechanism of the formation of bibenzyl from the reduction of benzyl bromides, and Fry and coworkers [221] described steric effects on the reductive coupling of α-substituted benzyl bromides. Other compounds that have been studied are benzhydryl halides [222,223] and α,α′dibromoxylenes [224]. Electrolyses of trifluoromethylbenzene, chlorodifluoromethylbenzene, and dichlorofluoromethylbenzene [225] at mercury afford a series of species, each possessing one less halogen atom than its precursor. Trifluoromethyl arenes undergo reductive coupling with electrophiles such as carbon dioxide, dimethylformamide, and acetone [226]. Trichloromethylbenzene undergoes stepwise reduction with the loss of chloride in DMF [103,227] to yield a mixture of 1,2-diphenyl-1,1,2,2-tetrachloroethane, diphenylacetylene, 1,2-dichloro-1,2-diphenylethene, and dichloromethylbenzene [227]; however, reduction of trichloromethylbenzene in the presence of acetic anhydride gives, besides the preceding products, 2-acetoxy-1-chloro-1-phenyl-1-propene and 1,2-diacetoxy-1-phenyl-1-propene [228]. Cathodic addition of trichloromethylbenzene to ketones gives α,β-unsaturated ketones [229]. Methyl esters of trifluoromethylbenzoic acid can be defluorinated by reduction at lead cathodes in MeOH [230,231]. Fry and Touster [232] electrolyzed a benzal halide in DMF containing trimethylchlorosilane to synthesize both the (α-halobenzyl)silane and the benzal disilane. Amino acids can be synthesized via the coupling of Schiff bases with carbanions generated by the reduction of benzyl halides [233,234]: PhCH2Cl

2e, –Cl– Hg cathode DMF, TBAI

PhCH2– CH3 1. PhCH2N

C

CO2C2H5

2. H+

CH3 H 2N

C

CH3 CO2H

CH2Ph

VI.

ALICyCLIC HALIDES

A.

MONOHALOGENATED SPECIES

H2 (Pd/C)

PhCH2NH

C

CO2C2H5

CH2Ph

An early concern about the reduction of substituted and optically active 1-bromocyclopropanes and 1-iodocyclopropanes at mercury and carbon cathodes was whether cleavage of the carbon–halogen bond involves a radical or carbanion intermediate [235–237]. Mann and Barnes [238] have discussed the mechanism of the reduction of the carbon–halogen bond in these species. Hazard and coworkers [239] investigated how the identity of the supporting electrolyte cation or the presence of an adsorbed alkaloid influences the stereochemistry of the reduction of derivatives of 1-bromo-1-carboxy2,2-diphenylcyclopropane and of optically active derivatives of 2,2-dibromocyclopropane [239–242]. Monobromocycloalkanes, ranging from cyclobutyl bromide to cyclohexadecyl bromide, have been examined polarographically in DMF [25,243]; and polarographic data for 1-bromo- and

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2-bromonorbornane [27,244–246], 1-bromobicyclo[2.2.2]octane [244,245], 1,7,7-trimethyl-3halobicyclo[2.2.1]heptan-2-ones [247], and derivatives of 1-bromo- and 2-bromo-7-oxabicyclo[2.2.1]heptane [246] have been compiled. Electrolyses of 1-bromo- and 1-iodonorbornane at mercury in DMF involve two-electron cleavage of the carbon–halogen bond to afford norbornane [248]; however, at a platinum cathode, the reductions of 1-bromo-, 1-chloro-, and 1-iodonorbornane give mainly norbornane, along with some 1,1′-binorbornyl [249]. In work involving the use of silver cathodes by Rondinini et al., the mechanism for the reduction of haloadamantanes was probed [250,251], and reductive coupling of a glycosyl halide (tri-O-acetyl-α-d-fucopyranosyl bromide) was studied [66,67,251,252]. Electrochemical reduction of 2-bromo or 2-chlorocycloalkanones leads to either coupling or coupling with ring contraction [253].

B.

DIHALOGENATED SPECIES

Electrolysis of 2,2-dichloronorbornane, 2-exo-bromo-2-endo-chloronorbornane, and 2-endobromo-2-exo-chloronorbornane at mercury cathodes in DMF is a two-electron process, giving a mixture of nortricyclene and endo-2-chloronorbornane [254]. Reduction of 7,7-dihalobicyclo[4.1.0]heptanes revealed that formation of the less stable endo-7-halobicyclo[4.1.0]heptane is preferred over the exo isomer [255,256]. According to Rifi [140,142], electrolysis of 1-bromo-3chlorocyclobutane at mercury yields bicyclo[1.1.0]butane, cyclobutane, and cyclobutene, and a similar pattern prevails for 1,3-dibromo-1,3-dimethylcyclobutane. Hoffmann and Voss [257] found that 1,3-dibromocyclopentane and 1,3-dibromocyclohexane undergo reductive ring-closure reactions at a platinum cathode to yield bicyclo[2.1.0]pentane and bicyclo[3.1.0]hexane, respectively. Low-temperature reduction of 1-bromo-4-chlorobicyclo[2.2.0]hexane at mercury in DMF affords Δ1,4-bicyclo[2.2.0]hexane, which can be trapped with cyclopentadiene [258,259]. Rifi [142], as well as Wiberg and coworkers [260], electrolyzed 1,4-dibromobicyclo[2.2.2]octane in DMF and obtained evidence for the transient formation of [2.2.2]propellane.  Rifi [142]  reported that electrolysis of 1,5-dibromobicyclo[3.2.1]octane yields the stable  [3.2.1]propellane, and Leibzon  and coworkers [261] found that reduction of 3,7-dibromo-3,7-dinitrobicyclo[3.3.1]nonane affords 3,7-dinitroadamantane in good yield. Adcock and coworkers [262] have studied the reductions of 1,3-dihaloadamantanes, 1,4-dihalobicyclo[2.2.2]octanes, and 1,3-dihalobicyclo[1.1.1]pentanes. For the reduction of 1,4-dihalonorbornanes at platinum in DMF, Wiberg and coworkers [249] found that the products are norbornane and 1,1′-binorbornyl. However, when 1,4-dihalonorbornanes are electrolyzed at mercury in DMF [263,264], the major products are norbornane and bis(1-norbornyl)mercury, along with a small amount of 1,1′-binorbornyl, and it was concluded that [2.2.1]propellane is an intermediate.

C. OTHER SPECIES Semmelhack and coworkers [265] reduced 1,2-dimethyl-4,5,6,6-tetrachlorospiro[2.3]hexadiene at mercury in DMF containing D2O to prepare 1,2-dimethyl-4,5,6-trichloro-6- deuteriospiro[2.3]hexadiene. Reductive monodechlorination of 4,4,8,8-tetrachlorodispiro[2.1.2.1]octane has been reported [266]. Selective trimethylsilylation of tetrachlorocyclopropene to afford 1-(trimethylsilyl)trichlorocyclopropene has been described [267]. Polyfluorocyclohexadienes can be reduced at a mercury cathode in aqueous EtOH [268]. Strelow and coworkers [269] have reviewed the literature on the reduction of 1,1,3,3-tetrachlorodispirocyclobutanes to 1,3-dichlorodispirobicyclo[1.1.0]butanes. Chlorinated insecticides (chlordane, aldrin, alodan, and endosulfan) can be electrochemically dechlorinated at lead cathodes in MeOH [270,271]. Lindane (γ-hexachlorocyclohexane) undergoes stepwise dechlorination at a palladium-modified carbon-cloth cathode in MeOH [272]; one-step dechlorination of lindane to afford benzene takes place at glassy carbon in DMF [273].

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VII. HALOgENATED AROMATIC COMPOUNDS A.

MONOHALOBENzENES

Originally, reduction of a monohalobenzene was believed to entail irreversible two-electron scission of the carbon–halogen bond followed by rapid protonation of the phenyl carbanion to give benzene. Andrieux and coworkers [15,274] provided quantitative evidence that in nonaqueous media a more intricate mechanistic picture for the reductive cleavage of these compounds must be drawn. As shown in reaction (25.1), a reversible, one-electron transfer occurs to give a radical–anion, which decomposes (reaction (25.2)) to yield the phenyl radical. Once formed, the phenyl radical can accept an electron (reaction (25.3)) or can react homogeneously with the initially generated radical–anion (reaction (25.4)), and the resulting phenyl anion can be protonated by the medium (reaction (25.5)): [C6H5X –]

C6H5X + e k

[C6H5X –]

C6H5

+ e

C6H5 + [C6H5X –] C6H5– +

H+

C6H5

+

(25.1) X–

(25.2)

C6H5–

(25.3)

C6H5– + C6H5X

(25.4)

C6H6

(25.5)

Other reactions to be considered are abstraction of a hydrogen atom from the solvent by the phenyl radical, homogeneous reduction of a solvent radical by the halobenzene radical–anion, and heterogeneous reduction of a solvent radical at the electrode surface. A paper by Isse and coworkers [275] provides a fresh look at dissociative electron-transfer processes for the reduction of aromatic bromides at carbon and silver electrodes. Examples of electrosyntheses based on the reduction of monohalobenzenes are numerous. Reductions of iodobenzene or bromobenzene at mercury [276] or bromobenzene at silver [277] in carbon dioxide–saturated DMF afford benzoic acid. Bromobenzene, 3-methoxybromobenzene, 4-bromofluorobenzene, 4-(trifluoromethyl)chlorobenzene, and 1,2,4-trichlorobenzene can be converted into aldehydes at a cadmium-coated stainless-steel cathode in DMF [278]. In the presence of trimethylchlorosilane, silylation products arise from the reduction of various aryl halides at platinum in DMF [165]. Reduction of p-iodoanisole at mercury in carbon dioxide– saturated DMF produces p-anisic acid, anisole, and bis(p-anisyl)mercury [279]. Electrolysis at a mercury cathode in diglyme containing a mixture of TBABF4 and dimethylpyrrolidinium tetrafluoroborate has been used by Kariv-Miller and Vajtner [280] to convert fluorobenzene to benzene. Chami and coworkers [281,282] electrolyzed substituted aryl halides in the presence of an olefin and a redox catalyst to synthesize arylated addition compounds; for example, the reduction of 4-chlorobenzonitrile at platinum in liquid ammonia containing 4-chlorostyrene and 4,4′-bipyridine affords 1-(p-chlorophenyl)-2-(p-cyanophenyl)ethane. Bromobenzene bearing an o-substituent can undergo electroreductive coupling with ethyl acrylate to form an intermediate used to prepare benzolactones [283–285]. Aromatic aldehydes can be electrosynthesized via the palladium-catalyzed carbonylation of aryl iodides (e.g., 4-iodophenol, iodobenzene, and 1-iodo-2-methylbenzene) in the presence of formic acid [286]. Rondinini and

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coworkers [250,251,287,288] have taken advantage of the electrocatalytic properties of silver for the reduction of aromatic halides (see also Chapter 24). Donnelly and coworkers [289,290] reported that the following intramolecular cyclization can be performed at a mild-steel cathode in MeCN: X

N

N

N

N N

N N

N Mild-steel cathode MeCN, TEABF4

F

F

Electrolysis of an aryl halide at a cadmium-modified nickel cathode in DMF induces a formylation reaction between aryl carbanions and the solvent [291]. Reduction of aryl halides gives an aryl carbanion which, by acting as a base to deprotonate a nitrile, can cause coupling of the nitrile with esters, aldehydes, and ketones [292,293]. Electrolysis of 1-iodonaphthalene at carbon in propylene carbonate produces an aryl carbanion that attacks the solvent to form an aryl ester [294], whereas reduction of aryl halides at copper, silver, and palladium cathodes in propylene carbonate can be used to prepare biaryl compounds [295]. Electrochemical trimethylsilylation of aryl halides can be effected at a stainless-steel or carbon-cloth cathode in THF–HMPA containing trimethylchlorosilane [296]. It has been shown [297–299] that aryl boronic acids and esters can be synthesized via the reductive coupling of aryl halides (such as p-chlorotoluene and o-bromoanisole) with a reagent such as a trialkylborate or pinacolborane; this work has been extended to the use of benzyl halides [300], polyhaloarenes [301], and allylic halides [302].

B. POLYHALOBENzENES Benzene bearing two different halogens (except for fluorine) can be reduced in a stepwise fashion. For example, electrolysis of p-bromoiodobenzene gives bromobenzene (which can be further reduced to benzene) [303], and reduction of p-fluoroiodobenzene yields fluorobenzene (which is not reducible) [304]. Polychlorobenzenes can be selectively dechlorinated in DMSO by electrolysis at mercury [305]. Studies have been made of the reductive dehalogenation of 1,3-difluorobenzene [280], the electrocarboxylation of dichlorobenzenes [306], the electrochemical behavior of the entire family of chlorinated benzenes [307] as well as their reductive carboxylation [308], and the reductive coupling between 1,4-dichlorobenzene and 2,6-di-tert-butyl phenoxide [309]. In a study of the reduction of di-, tri-, and tetrahalobenzenes at carbon in DMF [310], it was discovered that 1,2,4,5-tetrabromobenzene undergoes an electrolytically induced halogen dance. For the electrolysis of o-dihalobenzenes, benzyne has been proposed as an intermediate [311–315]. Mechanistic studies of the reductive coupling of polyhalogenated nitrobenzenes have been reported by Andrieux and coworkers [316]. A palladium-modified carbon-cloth cathode has been used by Kulikov and coworkers [272] for the stepwise dechlorination of 1,2,3,5-tetrachlorobenzene and by Tsyganok and Otsuka [317] to carry out the dechlorination of 2,4-dichlorophenoxyacetic acid. Electroreductive dechlorination of 5-chloro-2(2,4-dichlorophenoxy)phenol (triclosan) at a carbon cathode has been accomplished [318]. Voss and coworkers [319] employed lead cathodes for the dehalogenation of chlorinated dibenzofurans and dibenzo-p-dioxins; reductive carboxylation of the same compounds was investigated later [308].

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C.

Organic Electrochemistry

HALOGENATED NITROBENzENES AND CYANOBENzENES

Radical–anions formed by the addition of a single electron to a monohalobenzene are extremely short-lived. However, for some mono- and polyhalogenated nitrobenzenes and cyanobenzenes, the electrogenerated radical–anions are sufficiently stable that their EPR spectra can be recorded. Hawley [5] reviewed information about the electrochemistry of halogenated nitrobenzenes and cyanobenzenes, and Andrieux and coworkers [316] investigated the reduction of pentafluoronitrobenzene and pentachloronitrobenzene at glassy carbon in MeCN.

D.

HALOGENATED ARYL CARBONYLS

Product distributions have been determined for the electrolyses of pentafluorobenzaldehyde, pentafluorobenzoic acid, and pentafluorobenzamide at mercury in aqueous media [320]. Reduction of the aldehyde gives pentafluorobenzyl alcohol and 2,3,5,6-tetrafluorobenzyl alcohol. Pentafluorobenzoic acid is initially converted to a mixture of 2,3,5,6-tetrafluorobenzoic acid, pentafluorobenzyl alcohol, and 2,3,5,6-tetrafluorobenzyl alcohol; however, at a more negative potential, only the last two products are obtained. Pentafluorobenzaldehyde and 2,3,5,6-tetrafluorobenzaldehyde are easier to reduce than their benzoic acid analogues; therefore, the aldehydes cannot be isolated from electrolyses of the acids. In sulfuric acid, electrolysis of pentafluorobenzamide gives pentafluorobenzyl alcohol and 2,3,5,6-tetrafluorobenzyl alcohol. Using a lead cathode in sulfuric acid, Sato and coworkers [321] selectively reduced pentafluorobenzoic acid to either 2,3,5,6-tetrafluorobenzyl alcohol or 2,3,5,6-tetrafluorobenzaldehyde. Fry and coworkers [303] observed that electrolysis of m-bromoacetophenone at mercury in DMF affords acetophenone, whereas reduction of p-bromo-γ-chlorobutyrophenone gives γ-chlorobutyrophenone. Halogenated benzophenones behave in two different ways in nonaqueous solvents such as DMF, DMSO, and MeCN [322]. For m-fluoro-, p-fluoro-, and m-chlorobenzophenone, a pair of electrons is apparently involved in the reduction, and benzophenone is produced quantitatively. However, for the electrolysis of p-chloro-, m-bromo-, and p-bromobenzophenone, fewer than two electrons per molecule are transferred, and benzophenone is obtained in less than 100% yield. Reduction and carboxylation of halobenzophenones have been investigated by Isse and coworkers [323]. Reduction of 2-bromo- and 2-iodo-4′-methoxy-N-methylbenzanilide at mercury in DMF involves one-electron cleavage of the carbon–halogen bond to yield a radical intermediate [324,325], which can undergo (1) further reduction and protonation to give 4′-methoxy-N-methylbenzanilide, (2) intramolecular cyclization to afford 2-methoxy-N-methylphenanthridinone, or (3) rearrangement followed by hydrogen atom abstraction to yield 4′-methoxy-N-methylbiphenyl-2-carboxamide.

E.

HALOGENATED BIPHENYLS

Electroreduction of mono- and polychlorobiphenyls at mercury proceeds via sequential two-electron loss of chloride; however, due to the varied pattern of chlorine substitution, more than one product can form [326,327]. Maruyama and Murakami [328] reported that chlorinated biphenyls are reduced in DMF via one-electron cleavage of a carbon–chlorine bond to yield chloride and an aryl radical, which subsequently abstracts a hydrogen atom from the solvent. Several monochlorobiphenyls have been reduced at lead [307], and their reductive carboxylation has been achieved [308]. Rusling and Arena [329] have probed the reduction of 4-bromo-, 4,4′-dibromo-, 3,4-dichloro-, and 2,2′,5,5′-tetrachlorobiphenyl at mercury. Using a variety of electrodes, including glassy carbon, gold, palladium, and platinum, Simonet and Jouikov [330] found that reduction of 9-bromofluorene produces the 9-fluorenyl radical that appears to be grafted onto the surface of the cathode (see also Chapter 42).

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VIII.

ACyL HALIDES

Using mercury cathodes, Barba and coworkers have investigated the electrochemical reduction of benzoyl chloride [331], 1-naphthoyl chloride [331], 2-naphthoyl chloride [331], phthaloyl dichloride [332], cinnamoyl chloride [333], and 2- and 4-nitrobenzoyl chloride [334] (see also Chapter 32). Electrolysis of benzoyl chloride gives an acyl radical that dimerizes to yield benzil; the latter species accepts two electrons and undergoes double acylation by benzoyl chloride to afford a mixture of (E)- and (Z)-1,2-diphenylethene-1,2-diyl dibenzoate [232]. Cinnamoyl chloride [333] accepts one electron to form an acyl radical that abstracts a hydrogen atom from the medium to give an intermediate ketene, which then dimerizes. Reduction of benzoyl chloride and benzoyl fluoride at carbon and platinum cathodes in MeCN has been carried out by Cheek and Horine [335]. A study [336] of the effect of potential on the reduction of benzoyl chloride at mercury and carbon in MeCN revealed that acyl radicals and acyl anions are both involved as intermediates. In the same work [336], reduction of heptanoyl chloride was explored. Folest and coworkers [337] obtained symmetrical diketones by reducing aroyl and arylacetyl chlorides at nickel in MeCN. Other acyl chlorides that have been investigated include glutaryl dichloride [338], trimethylacetyl chloride [339], cyclohexanecarbonyl chloride [340], phenylacetyl chloride [341], hydrocinnamoyl chloride [341], phthaloyl dichloride [342], 2-furoyl chloride [343], and 2,4,6-trimethylbenzoyl chloride [344]. Quantitative conversion of 1-adamantanecarbonyl chloride to 1-adamantanecarboxaldehyde has been accomplished [345]. Lozano and Barba [346] reduced 2-chloro-2-phenylacetyl chloride at mercury to synthesize derivatives of pyran-2-one and pyran-4-one (see also Chapter 34): H Ph

C

O C

2e, –2Cl–

Ph

C H

C

O

Cl

Cl

Trimerization OH Ph

O Ph

Ph

Ph + O

Ph

O

O

OH

Ph

In addition, these workers investigated the reduction of 2-bromo-2,2-diphenylacetyl bromide at graphite in the presence of inorganic sulfur compounds and obtained a new sulfurated heterocyclic product [347–349].

IX.

PHENACyL HALIDES

Papers concerning the electrochemistry of phenacyl bromides have been published by Barba and coworkers [350–360]. Reduction of 2-bromo-1-phenylethanone at mercury affords 2,4-diphenylfuran [350,351]. Other compounds similarly synthesized include 2,4-bis(4-methoxyphenyl)furan, 2,4-bis(4-biphenylyl)furan, 2,4-bis(4-bromophenyl)furan, and 2,4-bis(4-nitrophenyl)furan. Products arising from the reduction of 2-bromo-1-phenylethanone are strongly influenced by the addition of proton donors and other reagents; for example, electrolysis of 2-bromo-1-phenylethanone in the presence of ethyl bromoacetate gives cyclopropane-1,2,3-triyltris(phenylmethanone) and 3,5-diphenylfuran-2(3H)-one [352].

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By protecting the carbonyl group of a phenacyl bromide through reaction with semicarbazide hydrochloride, one can synthesize a variety of 2,5-diarylfurans [353]. In addition, reductive dimerization of the semicarbazone derivative of 2-bromo-1-phenylethanone can be employed to form 3,6-diphenylpyridazine in excellent yield; this method is suitable for the preparation of a variety of 3,6-diarylpyridazines [354]. Reduction of α-bromopropiophenone at mercury gives 1,4-diphenyl-2,3-dimethylbutan-1,4-dione [355], but electrolysis of α-bromopropiophenone in the presence of benzoyl chloride affords only 1,3-diphenyl-2-methylpropan-1,3-dione. Other studies involving reduction of phenacyl bromides include the synthesis of 4-aryl-2-methylfurans [356] and of enol carbonates [357]. Semicarbazones of 2-bromo-1-phenylethanone can be converted into 3,7-diaryl-2H-imidazo[2,1-b][1,3,4]oxadiazines [358]. Reduction of 1,2-dibenzoylchloroethane at mercury in DMF affords mixtures of phenyl tribenzoyl cyclopentanols and diphenyl dibenzoyl butanediones [359]. Electrolysis of 2-bromo3-oxo-3-phenylpropanenitrile results in the formation of (2-amino-5-isocyano-4-phenylfuran-3-yl)(phenyl)methanone [360]: O NC O H5C6

C

NH2

Br CH CN

H5C6

C6H5 O

Electrolytic reduction of 2,2-dibromo-1-phenylethanone in DMF gives (E)-(3-(dibromomethyl)3-phenyloxiran-2-yl)(phenyl)methanone and trans-cyclopropane-1,2,3-triyl(phenylmethanone) [361].

X. ALIPHATIC α-HALOCARbONyL COMPOUNDS Polarographic studies of aliphatic α-haloaldehydes in water–dioxane [362,363] have shown that reduction of these compounds is influenced by the equilibrium involving the hydrated and unhydrated forms of the aldehyde. Each carbon–halogen bond is cleaved in a two-electron, one-proton process, and the final nonhalogenated aldehyde is reduced at a more negative potential [364]. At a mercury cathode in DMF, 2-bromo-1-phenylethanone can undergo either simple carbon–bromine bond cleavage or reductive coupling [365]. Reduction of 2-bromo-3-pentanone in DMF–H 2O affords a mixture of 3-pentanone and 1-hydroxy-3-pentanone [366]. In the same study, electrolysis of α,α′-dibromoacetone in the presence of benzoate resulted in both cleavage of a carbon–bromine bond and SN2 displacement of bromide by benzoate. In an acetic acid–sodium acetate buffer, reduction of branched dibromo ketones, such as 2,4-dibromo2,4- dimethyl-3-pentanone, gives α-acetoxyketones in excellent yield; less highly substituted species, such as 4,6-dibromo-5-nonanone, undergo reductive cleavage of both carbon–bromine bonds [367–369]. Depending on experimental conditions, reduction of bis(α-bromocyclopropyl) ketone in an acetic acid–sodium acetate buffer can yield either α-bromocyclopropyl cyclopropyl ketone or dicyclopropyl ketone [370]. Electrolyses of 2,4-dibromo-2-methylpentan-3-one and of 2,4-dibromo-1,5-diphenyl1,4-pentadien-3-one have been reported [371,372]. Low-temperature reduction of alkylated α,α′dihaloketones to cyclopropanones can be accomplished with highly alkylated starting materials, and electrochemical cyclization of 1,3-dihaloketones (with carbonyl moieties protected as acetals, acylals, or aminals) can produce cyclopropanone derivatives [373]. Using mercury cathodes in

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DMF, Inesi et al. [374] reduced α,α′-dichloroketones in the presence of carboxylic acids, phenols, or diethyl malonate to obtain a variety of products. Reduction of 1,1,3-tribromo- or 1,1,3-trichloro3-methyl-2-pentanone in the presence of amines or phenols gives the α,β-unsaturated amides or esters [375]: O

O H5C2

C C Cl

CH3

H C

Cl

C

H5C2 2e, –2Cl–

C

CH

H3C

Cl

Cl

PhCH2NH2

O

H5C2

C

C H3C

C H

CH2Ph N H

Inesi and coworkers [376] carried out the diastereoselective electrocarboxylation of chiral N-(2bromoacyl)oxazolidin-2-ones. In earlier work, reduction of α,α′-di- and trichloroketones in the presence of oxazolidin-2-ones provided a synthetic route to N-enoyloxazolidin-2-ones [377]. Barba and coworkers [378] reported that α-bromodibenzoylmethane is the only product obtained via reduction of α,α-dibromodibenzoylmethane at platinum or mercury in DMF. In aqueous media, α-halocarboxylic acids as well as their esters undergo two-electron, oneproton reductive cleavage [379–381]; reduction of chloro- and dichloroacetic acids and their ethyl esters has been studied in DMF [382]. Electrolysis of ethyl bromoacetate at mercury in slightly wet DMF gives mainly ethyl acetate and diethyl succinate, whereas reduction of ethyl trichloroacetate in the presence of acrylonitrile affords 1-carbethoxy-1-chloro-2-cyanocyclopropane and 1-carbethoxy-1,1,3-trichloro-3-cyanopropane [95]. Electroreductive activation of methyl 2,2,2-trichloroacetate in the presence of cyclohex-2-enone yields a bicyclic chlorocyclopropane [383]. Reduction of ethyl 2-bromo-2-phenylacetate affords a mixture of meso- and d,l-diethyl-2,3-diphenylsuccinate [384]; analogous behavior is exhibited by methyl 2-bromo2-phenylpropanoate [385] and by ethyl α-bromonaphthalene-1-acetate [386]. Direct and indirect reduction of ethyl 3-bromopropanoate and ethyl 4-bromopentanoate has been investigated with carbon, gold, and mercury cathodes [387]. Reactions of carbanions (formed via reduction of bromoesters) with β-, γ- and δ-bromoesters have been probed [388,389]. Electrolysis of diethyl 2,2′-(1,2-phenylene)bis(2-bromoacetate) affords a mixture of carbocyclic products [390]. Reduction of ethyl α-bromo-9-anthrylacetate at glassy carbon yields 9,10-anthraquinone [391]. Electrolysis of decyl dichloroacetate or trichloroacetate at stainless steel in the presence of a trialkylborane causes loss of a chlorine moiety and its replacement with an alkyl group [392]. Preparative-scale reduction of ethyl 2-bromo-3-(3′,4′-dimethoxyphenyl)-3-(propargyloxy)propanoate has been carried out at vitreous carbon cathodes [393]. Giomini et al. [394,395] investigated the reduction of ethyl α,β-dibromopropanoate, ethyl α,γ-dibromobutyrate, and methyl α,δ-dibromovalerate. Reduction of 2,2-dibromo-1H-indene-1,3(2H)-dione in dichloromethane

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has been studied [396]. Two papers deal with the electroreductive intramolecular cyclization of dimethyl dibromoalkanedioates [397,398]: O Br

O

OCH3

CH

H C

H3CO

OCH3

2e Pt cathode

OCH3

O

Br

O

Reduction of α-haloamides at mercury in DMF has been reported [399,400]; biologically active N-benzyloxazolidine-2,4-dione was prepared by the electrolysis of a solution containing α-chloroN-benzylacetamide, carbon dioxide, and a probase. Maran et  al. examined the electrochemical reductions of 2-bromocarboxamides [401] and 2-bromo-2-methylpropanamides [402] as well as the electrocarboxylation of 2-bromoisobutyramides [403]. In other research, Maran [404] studied the reduction of several α-haloamides, and Casadei and coworkers carried out the reductive dehalogenations of 1-acetamido-2,2,2-trichloroethyl acetate [405] and of substituted N-(2,2,2-trichloroethyl)acetamides [406–408]. Reduction of compounds such as 2-bromo-N-cyclohexylacetamide to N-cyclohexylacetamide has been examined [409], and reductive cyclization of diethyl 2-(2-bromoN-(4-methoxyphenyl)acetamido)malonate and its analogues affords a lactam [410,411]; reduction of a 3-halo-β-lactam in the presence of acetic anhydride gives the corresponding 3-acetyl-β-lactam [412]. Electrochemical reduction of additional haloamides has been carried out [413–416].

XI. HALOgENATED HETEROCyCLIC COMPOUNDS Halogenated heterocyclic substances tend to display electrochemical behavior similar to that of aryl halides; the carbon–halogen bond is reductively cleaved in a process involving two electrons and a proton to give a hydrogenated product.

A.

HETEROCYCLIC COMPOUNDS WITH ONE NITROGEN ATOM

Polarographic studies of aqueous solutions of monohalogenated pyridines (e.g., 4-chloro, 2-bromo-, and 3-iodopyridine), in either their unprotonated or protonated form, reveal a two-electron cleavage of the carbon–halogen bond, followed by protonation to give pyridine [417–422]. Reduction of 2-chloro-3-nitropyridine in DMSO initially yields a stable radical–anion [423], which is dechlorinated at a more negative potential to form the 3-nitropyridine radical–anion. Andrieux and coworkers [274] employed redox catalysis to probe the reduction of 2-bromo-, 3-bromo-, 2-chloro-, and 3-chloropyridine at mercury in DMF. Reduction of 2,5-dibromo-, 2,3-dichloro-, and 2,6-dichloropyridine in EtOH–water was studied by means of polarography [424], and bulk reduction of mono- and dihalopyridines at carbon cathodes in DMF has been carried out [425]. Reductive coupling of monohalopyridines with carbon dioxide affords the corresponding carboxylic acids [426]. Electrochemical reduction of pentafluoropyridine and pentachloropyridine has been reported [427], and several patents pertaining to reductive dechlorination of pentachloropyridine have appeared [428–430]. Gennaro and coworkers [431] synthesized 6-aminonicotinic acid by electrolyzing 2-amino-5-halopyridines at silver in the presence of carbon dioxide: CO2H

Br e, –Br–, e, CO2, H+ Ag cathode H2N

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N

H2N

N

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Alwair and Grimshaw [432] studied the reduction of 6-chloroquinoline. In subsequent research by Fuchs and coworkers [433], the electrochemical carboxylation of a number of chlorinated quinolines was investigated in the absence and presence of carbon dioxide. According to Kyriacou et al. [434], reduction of 3,4,5,6-tetrachloro-2-picolinic acid at a silver electrode in an aqueous medium yields 3,4,6-trichloro-2-picolinic acid, which can be further reduced to 3,6-dichloro-2picolinic acid.

B.

HETEROCYCLIC COMPOUNDS WITH TWO NITROGEN ATOMS

Electrochemical reduction of 6-chloroquinoxaline in DMF affords a stable radical–anion, whereas the bromo and iodo analogues accept one electron to yield radical–anions, which undergo fragmentation and hydrogen atom abstraction to form quinoxaline [432]. In similar fashion, 2-bromo- and 2-chlorophenazine are reversibly reduced to stable radical–anions, but the radical– anion of 2-iodophenazine decomposes to yield iodide ion and phenazine [432]. Lund [435] observed that reduction of 4-chloroquinazoline results in two-electron cleavage of the carbon–chlorine bond to form quinazoline. In addition, the electrochemical reduction of halogenated pyrimidines has been described in several publications [436–440]. A survey of the electrochemistry of chlorinated pyrazines, quinoxalines, and pyridazines has appeared [441].

C. OTHER HETEROCYCLIC SPECIES Laviron [417] investigated the polarography of halogenated thiazoles, and Iversen [442] synthesized thiazole in 88% yield by reducing 2-bromothiazole at mercury in an acidic water–EtOH medium. Reduction of 2-bromothiazole at carbon in MeCN involves both radical and carbanion intermediates, with thiazole being the only product [443]. Polarographic studies of brominated and iodinated thiophenes in aqueous and nonaqueous media have been undertaken [444–448]. Feldmann and Koberstein [449] probed the reduction of 2,3,5-tribromothiophene at mercury, and Pletcher and Razaq [450] synthesized 3-bromothiophene by reducing 2,3,5-tribromothiophene at various cathodes in dioxane–water. Reduction of tetrabromothiophene, tetraiodothiophene, 2,3,4-tribromothiophene, and 3,4-dibromothiophene in DMF has been described by Gedye and coworkers [451]. Electrosyntheses of 3-bromo- or 3,4-dibromothiophene, of 3,4-dibromo-2-chlorothiophene, and of 3-bromo-4-chlorothiophene have been developed by Dapperheld and coworkers [452]. Reduction of several dihalothiophenes at carbon in DMF involves an electrolytically induced halogen dance [453]. Justice and Hall [454] synthesized 3-hydrobenzo[b]thiophene by reducing a substituted 3-chlorobenzo[b]thiophene. Reduction of 3,3-dichloroisobenzofuran-1(3H)-one provides a pathway to the synthesis of 3-substituted phthalides [455]. Casadei and coworkers [456] studied the reduction of 3-halo-β-lactams and 3,3-dihalo-β-lactams in DMF. In the presence of carbon dioxide, a trans-3-halo-1,4-diphenylazetidin-2-one can be converted into trans-3-carboxy-1,4-diphenylazetidin-2-one; without carbon dioxide, the major product is the dehalogenated β-lactam.

XII. INDIRECT OR CATALyTIC CLEAVAgE OF CARbON–HALOgEN bONDS Direct electrolytic cleavage of a carbon–halogen bond is an intrinsically irreversible process. Thus, the potential needed to reduce a carbon–halogen bond heterogeneously at an electrode is often much more negative than the thermodynamically reversible potential. However, carbon–halogen bonds can be cleaved chemically with electron-transfer mediators (catalysts) that are electrogenerated at potentials more positive than those required for direct reduction of the carbon–halogen bond. In a classic application of this concept, Lund and coworkers [457] used the electrogenerated radical– anion of chrysene to reduce bromobenzene catalytically.

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Organic Electrochemistry

ELECTROGENERATED RADICAL–ANIONS AS MEDIATORS

Numerous studies [457–463] have shown that the rate of electron transfer between an electrogenerated radical–anion (mediator) and a halogenated organic compound (substrate) increases as the difference between the standard potentials for the reduction of the precursor of the mediator and for the reduction of the substrate decreases. Reactions of 1-bromo- and 1-chlorobutane with the electrogenerated radical–anions of trans-stilbene and anthracene in DMSO have been examined [458]. Cyclization of N-allyl-2chloroacetanilide to 1-(3-methylindolin-1-yl)ethanone occurs in the presence of the radical–anion of trans-stilbene [459]: O

O CH3

ClCH2

N

Radical-anion of trans-stilbene

N

CH3

For the reaction between alkyl halides and the electrogenerated naphthalene radical–anion, Sease and Reed [460] observed that only alkyl chlorides, such as 1-chlorohexane and 6-chloro-1-hexene, are catalytically reduced. Britton and Fry [461] elucidated the kinetics of the electron-transfer reaction between 1-chlorooctane and the phenanthrene radical–anion in DMF. Another feature of reactions between alkyl halides and electrogenerated aromatic radical–anions should be mentioned. After a single electron has been transferred to the alkyl halide (substrate) to form an alkyl radical, that radical can (1) react with another radical–anion (mediator) to form an alkyl carbanion (R−), which is protonated to form RH or (2) bond with the radical–anion to yield an alkylated aromatic hydrocarbon. An example of the latter process is seen in the electrolysis of stilbene in DMF containing tert-butyl chloride [464]. A recent expansion of this concept is the catalytic reduction of alkyl halides at silver–palladium cathodes in the presence of arenes to afford alkylated products [465]. Other studies of the reactions of alkyl halides with electrogenerated radical–anions include reduction of diketones in the presence of methyl halides [466], reduction of nitrobenzene in the presence of alkyl halides [467], reductive coupling of 1,2- and 1,3-dihaloalkanes with anthracene [468], reductive coupling of phenylacetylene with alkyl halides [469], electron transfer between meso- or d,l-1,2-dichloro-1,2-diphenylethane and electrogenerated radical–anions [470], and reductive coupling of tert-butyl bromide with azobenzene, quinoxaline, and anthracene [471]. Reactions between halogenated organic compounds and electrogenerated radical–anions have been extended to unsubstituted and substituted benzyl chloride [472]; to vinyl bromides, 2-bromoindene, and 7,7-dichlorobicyclo[4.1.0]heptane [473]; to bornyl, isobornyl, and exo- and endo-norbornyl bromides [474]; and to lindane [475]. Inesi [476] has investigated the indirect reduction of bromo esters by the electrogenerated radical–anions of pyrene, anthracene, diphenylanthracene, and fluoranthene.

B.

TRANSITION-METAL COMPLEXES AS MEDIATORS

Nédélec and coworkers [477] have reviewed the subject of organic electroreductive coupling reactions that utilize transition-metal complexes as catalysts. Durandetti and Périchon [478] published a short review of nickel-catalyzed coupling reactions involving the reactions of aryl, heteroaryl,

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and vinyl halides with activated alkyl halides. Papers by Duñach et al. [479,480] describe the use of nickel and palladium species as catalysts for the reduction of halogenated organic compounds. Classically, the two most important transition-metal catalysts for the reduction of halogenated organic compounds have been electrogenerated low-valent nickel and cobalt species. Papers by Pletcher and coworkers [481–485] introduced the catalytic reduction of various bromoalkanes by electrogenerated nickel(I) complexes; ligands coordinated to nickel(I) included Schiff bases and macrocyclic tetraamines. Although organonickel(III) intermediates were proposed for these reactions, no definitive evidence for these species has been reported. In contrast to the behavior of nickel(I) complexes as mediators, the catalytic reactions of alkyl halides with electrogenerated cobalt(I) species such as vitamin B12s, cobaloximes(I), and cobalt(I) salen exhibit a significant difference. Cobalt(I) species, acting as potent nucleophiles in SN2 reactions with alkyl halides, give stable and identifiable organocobalt(III) intermediates. Lexa and coworkers [486] articulated the mechanism for the catalytic reduction of 1-bromobutane by electrogenerated cobalt(I) tetraphenylporphin. In Sections XII.B.1 through XII.B.3, examples of catalytic reductions of halogenated organic compounds by electrogenerated nickel(0), nickel(I), cobalt(I), and other species are mentioned. 1. Nickel Complexes Electrogenerated (2,2′-bipyridine)nickel(0) complexes serve as catalysts for many syntheses: (1)  formation of ketones from acyl halides and either alkyl or aryl halides [487,488]; (2) production of β,γ-unsaturated esters via coupling of α-haloesters with aryl or vinyl halides [489]; (3) coupling of α-chloroesters or α-chloronitriles with carbonyl compounds [490]; (4) formation of unsymmetrical biaryls from a mixture of two aryl halides [491]; (5) reduction of alkyl halides or α,ω-dibromoalkanes to form dimers [492]; (6) coupling of alkyl, benzyl, and aryl halides [493]; (7)  formation of α-arylated ketones via coupling of α-chloroketones with aryl  halides  [494]; (8)  cross-coupling of aryl halides and ethyl chloroacetate [495]; (9) coupling of aryl halides and activated olefins [496]; (10) cross-coupling between aryl halides and activated alkyl halides [497]; (11) formation of ketones via coupling of organic halides with carbon monoxide [498–500]; (12) production of ketones via reduction of a mixture of a benzyl or alkyl halide with a metal carbonyl [501]; (13) preparation of symmetrical ketones from alkyl, benzyl, and aryl halides [502]; (14) synthesis of 2-arylpyrimidines and 2-arylpyrazines from 2-chloropyrimidine and 2-chloropyrazine (in the presence of aryl halides) [503]; (15) formation of aryl thioethers from aryl halides and thiophenol [504]; (16) preparation of aryl propan-2-ones via coupling of benzylic chlorides with an acyl donor species [505]; (17) dehalogenation of chlorinated benzenes and dibenzofurans [506]; (18) homocoupling of 2-bromomethylpyridines [507]; (19) electrocarboxylation of substituted aryl halides [508]; (20) synthesis of 2-arylpyridines via cross-coupling of aryl and pyridyl halides [509,510]; (21) coupling of 2- and 3-bromothiophene with alkyl and alkenyl halides [511]; (22) synthesis of 3-thienylzinc bromide from 3-bromothiophene [512]; (23) preparation of unsymmetrical ketones via reduction of mixtures of benzyl and aryl halides in the presence of iron pentacarbonyl [513]; (24) synthesis of 2,2-difluoro-3-hydroxyesters from methyl chlorodifluoroacetate and carbonyl compounds [514]; (25) preparation of conjugated dienes via homocoupling of alkenyl halides [515]; and (26) coupling of aryl halides with chlorodiphenylphosphine or dichlorophenylphosphine [516]. Additional electrosyntheses based on the use of (2,2′-bipyridine)nickel(0) species include crosscoupling of aryl halides with arenecarboxaldehydes to afford arylated secondary alcohols [517,518], reductive polymerization of 3-substituted 2,5-dihalothiophenes [519], homocoupling of 2-bromomethylpyridines [520], cross-coupling of aryl halides with 3-chloropyridazines to yield arylpyridazines [521], homo- and cross-couplings of alkenyl halides [522], cross-coupling between aryl halides and α-chloropropanoic acids bearing chiral auxiliaries [523], synthesis of arylzinc compounds via reduction of aryl halides in a cell with a sacrificial zinc anode [524], reductive coupling

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of a variety of organic halides [525,526], coupling of methyl 2-chloroacetate with ketones [173], coupling of 3-chloro-2-methylprop-1-ene with ketones and aldehydes [173], homocoupling of o-substituted aryl halides and their heterocoupling with m- or p-substituted aryl halides [527], coupling of allylic chlorides or acetates with carbonyl compounds to form homoallylic alcohols [528], and coupling of α-haloesters with aryl or vinyl halides [489,529]. Nickel-catalyzed formation of cyclopropyl derivatives via coupling of activated olefins with gem-dibromoalkanes has been described [530], along with homocoupling of alkyl halides as well as cross-coupling of organic halides with activated olefins in an ionic liquid [531]. Reductive nickelcatalyzed conjugate addition of (Z)- or (E)-alkenyl halides to electron-deficient olefins affords functionalized pure (Z)- or (E)-olefins in high yield [532], and heteroaryl halides (e.g., monohalogenated pyridines, quinolines, and thiophenes) can be coupled with 3-buten-2-one [533]. Nickel(0) in the presence of potassium iodide promotes the conversion of nonactivated aryl bromides to aryl iodides [534]. Dimerization of halobenzenes as well as arylation of activated olefins can be accomplished with electrogenerated low-valent nickel complexes involving bipyridylamine ligands [535]. Electrogenerated nickel(0) pyridine complexes serve as catalysts for the coupling of aryl halides with 3,3-diethoxyprop-1-ene [536], the conjugate addition of substituted aryl bromides to but-3-en-2-one [537], and the synthesis of lactones via the arylation of α,β-unsaturated carboxylic esters [538]. Nickel(0) triphenylphosphine species have been used for the conversion of allyl halides to 1,5-hexadiene [539], for the reductive coupling of ethylene with aryl halides to give 1,1-diarylethanes [540,541] and styrene [542], for the coupling of aryl halides and alkenes to prepare substituted olefins [543], for the synthesis of ethyl 2-phenylacetate from ethyl-2-haloacetates [544], for the preparation of biaryls from aryl halides [545,546], and for the electrosynthesis of aryl carboxylates from aryl halides and carbon dioxide [547]. Nickel(0) tributylphosphine complexes have been used to catalyze the reduction of aryl and alkyl halides [548]. In addition, 1,2-bis[(di-2-propylphosphino)benzene]nickel(0) has been used for the reductive coupling of aryl halides [549]; 1,3-bis[(diphenylphosphino)propane]nickel(0) catalyzes the reductive carboxylation of 1-chloroethylbenzenes [550]; and 1,2-bis[(diphenylphosphino)ethane]nickel(0) has been employed to synthesize biphenyl from bromobenzene [551,552], to prepare benzoic acid from bromobenzene in the presence of carbon dioxide [553,554], to electropolymerize 1,4-dibromobenzene into poly(1,4-phenylene) [555], and to catalyze the reductive carboxylation of bromobenzene to benzoic acid [556]. Electrogenerated nickel(I) salen has been used as a catalyst for the reductive intramolecular cyclizations of 6-bromo- and 6-iodo-1-phenyl-1-hexyne [557], 6-bromo-1-hexene [558–561], and other acetylenic halides [562]; the reduction of several α,ω-dihaloalkanes [563]; the reductive coupling of 2-bromo- and 2-iodoethanol to prepare 1,4-butanediol [564]; the conversion of cyclohexanecarbonyl chloride to a tetramer [565]; electrocarboxylation of benzylic halides [566] and arylethyl chlorides [567]; and the reductions of benzal chloride [568] and 1-bromooctane [569]. In the presence of dioxygen, light, and water, electrogenerated nickel(I) salen can promote the reduction of alkyl monohalides to aldehydes and ketones [570–572]. An unanticipated consequence of the use of nickel(I) salen (or of nickel(I) with salen-like ligands) is that the phenyl-conjugated imino bonds of the ligand become alkylated by fragments of the original substrate [573,574]; this phenomenon, which destroys the action of the nickel(I) species as a mediator, is linked to the occurrence of ligand-centered (instead of metal-centered) reduction of the parent nickel(II) complex [575], a problem that can be overcome, at least in part, by the synthesis of ligands with sterically bulky groups on the imino bonds [576]. Electrogenerated nickel(I) salen has been used to promote ring-expansion reactions of some 1-haloalkyl-2oxocycloalkanecarboxylates [577]. Duñach and coworkers have utilized electrogenerated nickel(I) salen species for the intramolecular cyclizaton of some unsaturated 2-bromophenyl ethers [578,579]:

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Aliphatic and Aromatic Halides Br CH3 O

CH3

Nickel(I) C cathode DMF, TBABF4 H3C

H3C

CH2

CH3 +

O

O

Various electrogenerated nickel(I) tetraazamacrocycles with ligands such as cyclam, tetramethylcyclam, and hexamethylcyclam act as catalysts for the reduction of allyl o-halophenyl ethers [580] as well as the reductive intramolecular cyclizations of o-haloaryl compounds possessing unsaturated side chains [581–583] and of bromo propargyloxy esters [584–589]. Recent papers have extended the preceding applications to environmentally friendly media such as alcohols and water–alcohol mixtures [590–595]. Espenson and coworkers used such catalysts to reduce alkyl and benzyl halides [596– 599], α,ω-dihaloalkanes [600,601], 1-bromo-4-cyanobutane [602], and other halogenated compounds [603]. Using this same class of catalysts, Ozaki and coworkers carried out a number of interesting electrosyntheses: (1) intramolecular cyclization of 2-bromoalkyl- and 3-bromoalkyl-2-cyclohexen-1ones [604]; (2) radical cyclization of halogenated ethers [605]; (3) intermolecular addition of alkyl radicals to activated olefins [606]; (4) intramolecular cyclization of n-allylic and n-propargylic α-bromoamides and of o-bromoacryloylanilides to give five-membered lactams [607]; (5) cyclization of vinyl and aryl halides [608]; (6) cyclization of α-bromo- and α-iodoamides [609]; (7) cyclization of acetylenic halides to prepare functionalized (methylene)cyclopentanes [610]; (8) stereoselective addition of n-, sec-, and tert-butyl radicals to α-methylenebutyrolactones [611]; (9) reduction of 2-bromoand 2-iodo-1,6-dienes to give bicyclo[3.1.0]-, 3-azabicyclo[3.1.0]-, and 3-oxabicyclo[3.1.0]hexane derivatives [612], and preparation of pyrrolopyridines and pyrrolopyrrole derivatives via reductive cycloaddition of 1-(2-iodoethyl)pyrrole to activated olefins or cyclization of 1-(ω-iodoalkyl)pyrroles [613]. Production of ethylene oxide via catalytic reduction of 2-haloethanol by nickel(I) cyclam has been reported [614]. Gómez and coworkers [615] converted nickel and palladium tetraazamacrocyclic complexes to their corresponding zero-valent states to reduce unsaturated o-haloaryl and o-halobenzyl ethers; the nickel(0) complex induced intramolecular cyclization, whereas the palladium(0) complex caused cleavage of the carbon–oxygen bond of the ethers. Electrogenerated nickel(I) cyclam has been used by Pelletier et al. [616] to produce dihydrobenzo[b]thiophenes from o-haloaryl allyl thioethers and by Nunnecke and Voss [506] to dehalogenate chlorinated benzenes and dibenzofurans. Stolzenberg and coworkers used electrogenerated nickel(I) tetrapyrrole complexes for the catalytic reduction of dichloromethane and methyl iodide [617], alkyl halides [618–620], and aryl halides [620]. Catalytic reduction of trans-1,2-dibromocyclohexane to cyclohexene by electrogenerated nickel(I), cobalt(I), and iron(I) porphyrin complexes was investigated by Lexa et al. [621], whereas reactions of simple alkyl bromides with aromatic radical–anions and low-valent iron porphyrins were considered in a later study [622]. Electrochemical polymerization of 1,4-bis(chloromethyl) benzene in the presence of catalytic amounts of nickel(II) chloride has been employed to prepare poly-p-xylylene [623]. Functionalized indanes and naphthalenes have been electrosynthesized via a nickel-catalyzed arylation of activated olefins (e.g., acrylonitrile or ethyl acrylate) [624]. 2. Cobalt Complexes Allyl halides have been reduced with electrogenerated tris(2,2′-bipyridine)cobalt(I) to afford 1,5-hexadiene [625,626]. Early work with cobalt(I) salen involved its use for the catalytic reduction

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of bromoethane [627], bromobenzene [627], and tert-butyl bromide and chloride [628]. Fry and coworkers examined the cobalt(I) salen-catalyzed reductions of benzal chloride [629–631] and of benzotrichloride [632]. Cobalt(I) salen-mediated reductions of 1-bromobutane [633,634], 1-iodobutane [634], 1,2-dibromobutane [634], benzyl and 4-(trifluoromethyl)benzyl chlorides [635], arylmethyl chlorides [636], iodoethane [637], diphenyl disulfide [638], 1,8-diiodooctane [639], 3-chloro-2,4-pentanedione [640], ethyl chloroacetate [641], 2,6-bis(chloromethyl)pyridine [642], (3-chloro-1-propen-1-yl)benzene [643], and 1-halonaphthalenes [644] have been investigated. Folest et  al. [645] studied the cobalt(I) salen-promoted carboxylation of benzylic and allylic chlorides. A  report concerning the catalytic reduction of ethyl chloroacetate by electrogenerated cobalt(I) salophen has appeared [646]. Rusling and coworkers carried out extensive studies of the use of electrogenerated cobalt(I) complexes (including cobalt(I) salen, vitamin B12s, and cobalt(I) phthalocyanine) as catalysts both in homogeneous phase and in bicontinuous microemulsions [647] (see also Chapter 8) for the reductions of 1,2-dibromoethane and 1,2-dibromobutane [648]; the debromination of alkyl vicinal dibromides [649]; the dechlorination of DDT [650]; the reductions of 1-bromobutane, 1-bromododecane, and trans-1,2-dibromocyclohexane [651–653]; the reduction of benzyl bromide [654]; the addition of primary alkyl iodides to 2-cyclohexen-1-one to afford 3-alkylcyclohexanones [655]; and the reductive intramolecular cyclization of 2-(4-bromobutyl)-2-cyclohexen-1-one to 1-decalone [653,655]. Cobalt(I) salen is a mediator for the dechlorination of DDT and its less chlorinated congeners [656], CFC-113 [657,658], CFC-113a [659], and hexa- and pentachlorobenzene [660]. Moreover, 4-methylcoumarin has been synthesized via the cobalt(I) salen-catalyzed reduction of 2-acetylphenyl 2-chloroacetate or 2-acetylphenyl 2,2-dichloroacetate [661]. Hisaeda and coworkers have synthesized a number of vitamin B12 model compounds with either one or two metal centers. Studies of the electrochemical behavior of these species, especially with respect to the use of their cobalt(I) states as catalysts for the reductive cleavage of carbon– halogen bonds, have been undertaken. Among the halogenated compounds that have been catalytically reduced are DDT as well as less-chlorinated members of the DDT family [662–664], (2-bromoethyl)benzene [665], benzyl bromide [666], diethyl 2-(bromomethyl)-2-methylmalonate [667–670], diethyl 2-(bromomethyl)-2-phenylmalonate [671], alkyl bromides bearing electronwithdrawing groups (e.g., 4-bromo-3-methyl-2-butanone, methyl 3-bromo-2-methylpropanoate, and 3-bromo-2-methylpropionitrile) [672,673], ethyl 4-bromo-5-oxohexanoate [674], 1-ethyl 3-phenyl 2-(bromomethyl)-2-methylmalonate [675], and a family of ethyl 1-(bromomethyl)-2oxocycloalkanecarboxylates [676]. Other uses of cobalt(I) catalysts include the reductive intramolecular cyclization of bromocyclohexenones to form bicyclic ketones [677] and the radical cyclization of bromoacetals [678,679]. Cross-coupling of functionalized aryl halides catalyzed by cobalt(I) pyridine complexes has been achieved [680]. Arylcobalt(II) complexes are reportedly formed via the reduction of aryl bromides in DMF–pyridine containing cobalt(II) bromide [681]. Electrogenerated cobalt(I) pyridine species can be used to cross-couple aryl halides with allylic esters [682,683]. Kräutler and coworkers [684] found that 1,4-dibromobutane interacts with electrogenerated cob(I)alamin to afford a tetramethylene-1,4-di-Coβ-cobalamin species. In a study of the reactions of cobalt(I) tetraphenyl porphyrin with benzyl chloride or 1-chlorobutane, Zheng and coworkers [685] reported that alkyl radicals are transferred from the cobalt center to a nitrogen of a pyrrole ring, leading to formation of an N-alkyl cobalt porphyrin complex. Electrogenerated (2,2′-bipyridine)cobalt(I) has been used to promote the Heck reaction between aryl halides and ethyl acrylate [686], to catalyze the addition of aryl halides to activated olefins [687], and to couple aryl halides with vinylic acetates to give styrenes [688–690]. Electrogenerated cobalt(I) complexes [691] can be employed to cross-couple aryl halides with 4-chloroquinolines [692], to synthesize arylzinc species from aryl halides [693–696], to produce organodizinc compounds from aromatic dihalides [697], and to couple aryl halides and allylic acetates [698].

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3. Other Catalysts Several papers have appeared that describe the use of electrogenerated samarium(II) for the catalytic reduction of organic halides at a nickel cathode in DMF [699], for the reductive coupling of 3-chloropropanoate with ketones to yield γ-butyrolactones [700], and for the reductive coupling of allyl chloride with ketones [701]. Torii et  al. [702] developed procedures for the palladium(0) triphenylphosphine-catalyzed carboxylation of aryl halides, β-bromostyrene, and allyl acetates. Homocoupling of aryl halides, catalyzed by anionic aryl-ligated palladium(0) bis(triphenylphosphine) species, has been examined [703]. Aromatic aldehydes have been electrosynthesized via the palladium-catalyzed carbonylation of aryl iodides in the presence of formic acid [286]. Hall and coworkers [704] electrogenerated tris(acetylacetonato)iron(II) for the reductive coupling and disproportionation of 1-bromooctane, and Wade and Castro [705] investigated the reactions of iron(II) porphyrin with vicinal dihalides as well as alkyl, allyl, benzyl, and propargyl halides. Catalytic dehalogenation of ethylene dibromide and trichloroacetic acid has been accomplished with iron(II) myoglobin in biomembrane-like surfactant films on pyrolytic graphite electrodes [706]. A report by Buriez et al. [707] concerns the mediation by iron(I) bipyridine complexes of the formation of β-hydroxyesters via coupling of α-haloesters and carbonyl compounds. Vanhoye and coworkers [708] synthesized aldehydes by using the electrogenerated radical–anion of iron pentacarbonyl to reduce iodoethane and benzyl bromide in the presence of carbon monoxide. Esters can be prepared catalytically from alkyl halides and alcohols in the presence of iron pentacarbonyl [709]. Yoshida and coworkers reduced mixtures of organic halides and iron pentacarbonyl and then introduced an electrophile to obtain carbonyl compounds [710] and converted alkyl halides into aldehydes by using iron pentacarbonyl as a catalyst [711,712]. A review by Torii [713] provides an overview of papers that deal with catalytic processes involving complexes of nickel, cobalt, iron, palladium, rhodium, platinum, chromium, molybdenum, tungsten, manganese, rhenium, tin, lead, zinc, mercury, and titanium.

C. REDUCTION OF CARBON–HALOGEN BONDS AT CHEMICALLY MODIFIED ELECTRODES Some papers have appeared that deal with the use of electrodes whose surfaces are modified with materials suitable for the catalytic reduction of halogenated organic compounds. Kerr and coworkers [714] employed a platinum electrode coated with poly-p-nitrostyrene for the catalytic reduction of 1,2-dibromo-1,2-diphenylethane. Catalytic reduction of 1,2-dibromo-1,2-diphenylethane, 1,2-dibromophenylethane, and 1,2-dibromopropane has been achieved with an electrode coated with covalently immobilized cobalt(II) or copper(II) tetraphenylporphyrin [715]. Carbon electrodes modified with meso-tetra(p-aminophenyl)porphyrinatoiron(III) can be used for the catalytic reduction of benzyl bromide, triphenylmethyl bromide, and hexachloroethane when the surface-bound porphyrin is in the Fe(I) state [716]. Metal phthalocyanine-containing films on pyrolytic graphite have been utilized for the catalytic reduction of trans-1,2-dibromocyclohexane and trichloroacetic acid [717], and copper and nickel phthalocyanines adsorbed onto carbon promote the catalytic reduction of 1,2-dibromobutane, trans-1,2-dibromocyclohexane, and trichloroacetic acid in bicontinuous microemulsions [718]. When carbon electrodes coated with anodically polymerized films of nickel(II) salen are cathodically polarized to generate nickel(I) sites, it is possible to carry out the catalytic reduction of iodoethane and 2-iodopropane [34] and the reductive intramolecular cyclizations of 1,3-dibromopropane and of 1,4-dibromo- and 1,4-diiodobutane [719]. Vaze and Rusling [720] employed a vitamin B12 analogue immobilized on carbon cloth for the catalytic reduction of 1,2-dibromocyclohexane, and Njue and Rusling [721] used a similar strategy for the catalytic reductive cyclization of n-bromoalkyl-2-cyclohexanones. A  volume edited by Murray [722] contains a valuable set of review chapters by experts in the field of chemically modified electrodes.

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Aliphatic and Aromatic Halides 570. 571. 572. 573. 574. 575. 576. 577. 578. 579. 580. 581. 582. 583. 584. 585. 586. 587. 588. 589. 590. 591. 592. 593. 594. 595. 596. 597. 598. 599. 600. 601. 602. 603. 604. 605. 606. 607. 608. 609. 610. 611. 612. 613. 614. 615. 616. 617.

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Aliphatic and Aromatic Halides 670. 671. 672. 673. 674. 675. 676. 677. 678. 679. 680. 681. 682. 683. 684. 685. 686. 687. 688. 689. 690. 691. 692. 693. 694. 695. 996. 697. 698. 699. 700. 701. 702. 703. 704. 705. 706. 707. 708. 709. 710. 711. 712. 713. 714. 715. 716. 717. 718. 719. 720. 721. 722.

979

Murakami, Y.; Hisaeda, Y.; Tashiro, T.; Matsuda, Y. Chem. Lett. 1985, 1813–1816. Tahara, K.; Chen, Y.; Pan, L.; Masuko, T.; Shimakoshi, H.; Hisaeda, Y. Chem. Lett. 2011, 40, 177–179. Murakami, Y.; Hisaeda, Y.; Fan, S. D.; Matsuda, Y. Chem. Lett. 1988, 835–838. Murakami, Y.; Hisaeda, Y. Pure Appl. Chem. 1988, 60, 1363–1368. Murakami, Y.; Hisaeda, Y.; Ozaki, T.; Matsuda, Y. J. Chem. Soc. Chem. Commun. 1989, 1094–1096. Murakami, Y.; Hisaeda, Y.; Ozaki, T. J. Coord. Chem. 1991, 23, 77–89. Hisaeda, Y.; Takenaka, J.; Murakami, Y. Electrochim. Acta 1997, 42, 2165–2172. Scheffold, R.; Dike, M.; Dike, S.; Herold, T.; Walder, L. J. Am. Chem. Soc. 1980, 102, 3642–3644. Torii, S.; Inokuchi, T.; Yukawa, T. J. Org. Chem. 1985, 50, 5875–5877. Begley, M. J.; Bhandal, H.; Hutchinson, J. H.; Pattenden, G. Tetrahedron Lett. 1987, 28, 1317–1320. Gomes, P.; Fillon, H.; Gosmini, C.; Labbé, E.; Périchon, J. Tetrahedron 2002, 58, 8417–8424. Buriez, O.; Kazmierski, I.; Périchon, J. J. Electroanal. Chem. 2002, 537, 119–123. Gomes, P.; Gosmini, C.; Périchon, J. J. Org. Chem. 2003, 68, 1142–1145. Gomes, P.; Gosmini, C.; Périchon, J. Org. Lett. 2003, 5, 1043–1045. Kräutler, B.; Dérer, T.; Liu, P.; Mühlecker, W.; Puchberger, M.; Gruber, K.; Kratky, C. Angew. Chem. Int. Ed. Engl. 1995, 34, 84–86. Zheng, G.; Stradiotto, M.; Li, L. J. Electroanal. Chem. 1998, 453, 79–88. Gomes, P.; Gosmini, C.; Nédélec, J. Y.; Périchon, J. Tetrahedron Lett. 2002, 43, 5901–5903. Gomes, P.; Gosmini, C.; Nédélec, J. Y.; Périchon, J. Tetrahedron Lett. 2000, 41, 3385–3388. Gomes, P.; Gosmini, C.; Périchon, J. Tetrahedron 2003, 59, 2999–3002. Polleux, L.; Labbé, E.; Buriez, O.; Périchon, J. Chem. Eur. J. 2005, 11, 4678–4686. Amatore, M.; Gosmini, C.; Périchon, J. Eur. J. Org. Chem. 2005, 989–992. Buriez, O.; Cannes, C.; Nédélec, J. Y.; Périchon, J. J. Electroanal. Chem. 2000, 495, 57–61. Le Gall, E.; Gosmini, C.; Nédélec, J. Y.; Périchon, J. Tetrahedron Lett. 2001, 42, 267–269. Gosmini, C.; Rollin, Y.; Nédélec, J. Y.; Périchon, J. J. Org. Chem. 2000, 65, 6024–6026. Buriez, O.; Nédélec, J. Y.; Périchon, J. J. Electroanal. Chem. 2001, 506, 162–169. Seka, S.; Buriez, O.; Nédélec, J. Y.; Périchon, J. Chem. Eur. J. 2002, 8, 2534–2538. Seka, S.; Buriez, O.; Périchon, J. Chem. Eur. J. 2003, 9, 3597–3603. Fillon, H.; Gosmini, C.; Nédélec, J. Y.; Périchon, J. Tetrahedron Lett. 2001, 42, 3843–3846. Gomes, P.; Buriez, O.; Labbé, E.; Gosmini, C.; Périchon, J. J. Electroanal. Chem. 2004, 562, 255–260. Hebri, H.; Duñach, E.; Périchon, J. Synth. Commun. 1991, 21, 2377–2382. Hebri, H.; Duñach, E.; Périchon, J. J. Chem. Soc. Chem. Commun. 1993, 499–500. Hebri, H.; Duñach, E.; Périchon, J. Tetrahedron Lett. 1993, 34, 1475–1478. Torii, S.; Tanaka, H.; Hamatani, T.; Morisaki, K.; Jutand, A.; Pluger, F.; Fauvarque, J. F. Chem. Lett. 1986, 169–172. Amatore, C.; Carre, E.; Jutand, A.; Tanaka, H.; Ren, Q.; Torii, S. Chem. Eur. J. 1996, 2, 957–966. Hall, J. L.; Geer, R. D.; Jennings, P. W. J. Org. Chem. 1978, 43, 4364–4366. Wade, R. S.; Castro, C. E. J. Am. Chem. Soc. 1973, 95, 226–230. Nassar, A. E. F.; Bobbitt, J. M.; Stuart, J. D.; Rusling, J. F. J. Am. Chem. Soc. 1995, 117, 10986–10993. Buriez, O.; Durandetti, M.; Périchon, J. J. Electroanal. Chem. 2005, 578, 63–70. Vanhoye, D.; Bedioui, F.; Montreux, A.; Petit, F. Tetrahedron Lett. 1988, 29, 6441–6442. Hashiba, S.; Fuchigami, T.; Nonaka, T. J. Org. Chem. 1989, 54, 2475–2476. Yoshida, K.; Kunugita, E.; Kobayashi, M.; Amano, S. Tetrahedron Lett. 1989, 30, 6371–6374. Yoshida, K.; Kobayashi, M.; Amano, S. J. Chem. Soc. Perkin Trans. 1 1992, 1127–1129. Yoshida, K.; Kuwata, H. J. Chem. Soc. Perkin Trans. 1 1996, 1873–1877. Torii, S. Synthesis 1986, 873–886. Kerr, J. B.; Miller, L. L.; Van De Mark, M. R. J. Am. Chem. Soc. 1980, 102, 3383–3390. Rocklin, R. D.; Murray, R. W. J. Phys. Chem. 1981, 85, 2104–2112. Elliott, C. M.; Marrese, C. A. J. Electroanal. Chem. 1981, 119, 395–401. Zhang, H.; Rusling, J. F. Talanta 1993, 40, 741–747. Kamau, G. N.; Rusling, J. F. Langmuir 1996, 12, 2645–2649. Dahm, C. E.; Peters, D. G. J. Electroanal. Chem. 1996, 406, 119–129. Vaze, A.; Rusling, J. F. J. Electrochem. Soc. 2002, 149, D193–D197. Njue, C. K.; Rusling, J. F. Electrochem. Commun. 2002, 4, 340–343. Murray, R. W., ed. Molecular Design of Electrode Surfaces, Techniques of Chemistry; Vol. XXII; Wiley: New York, 1992.

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26

Oxygen-Containing Compounds Alcohols, Ethers, and Phenols Robert Francke, Thomas Quell, Anton Wiebe, and Siegfried R. Waldvogel

CONTENTS I.

Aliphatic, Allylic, and Benzylic Alcohols ............................................................................ 982 A. Oxidation ...................................................................................................................... 982 1. Direct Oxidation .................................................................................................... 982 2. Indirect Oxidation .................................................................................................. 986 B. Reduction ......................................................................................................................990 1. Direct Reduction .................................................................................................... 991 2. Cathodic Deoxygenation after Introduction of Leaving Groups ........................... 991 II. Electrochemical Conversion of Ethers ................................................................................. 993 A. Oxidation ...................................................................................................................... 993 1. Saturated Ethers .....................................................................................................994 2. Unsaturated Ethers ................................................................................................. 995 B. Reduction ......................................................................................................................999 III. Aromatic Alcohols (Phenols)..............................................................................................1000 A. Oxidation ....................................................................................................................1000 1. Acetoxylation ....................................................................................................... 1001 2. Quinone Formation .............................................................................................. 1003 3. Phenol Coupling....................................................................................................1010 B. Reduction .....................................................................................................................1018 1. Electroreductive Hydroxylation of Aromatic Compounds ..................................1018 2. Electroreductive Birch-Type Reaction .................................................................1021 3. Electroreductive Hydrogenation and Hydrodeoxygenation of Phenolic Moieties ... 1024 4. Deoxygenation of Phenolic Compounds .............................................................1025 5. Reduction of Quinones ........................................................................................ 1026 References .................................................................................................................................... 1029

981

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982

I.

Organic Electrochemistry

ALIPHATIC, ALLyLIC, AND bENZyLIC ALCOHOLS

The hydroxy group represents the most abundant functional group and is therefore a very common moiety. Consequently, its electrochemical behavior in electroorganic conversion is of high interest. Due to the ability of alcohols to dissolve salts and their low reactivity, simple alcohols often serve as solvents or electrolyte components for electrochemical conversions.

A. OXIDATION 1. Direct Oxidation Generally, a direct electrochemical oxidation of alcohols is not very practical for preparative purposes. As high potentials have to be applied, most functional groups are not inert at the required conditions [1,2]. In particular, most aliphatic alcohols are unsuitable for direct electrolysis due to their extremely high stability toward anodic oxidation (see Table 26.1). Usually, the free electron pairs of a hydroxy group are more difficult to ionize than π electron systems of aromatic systems, which can be observed using photoelectron spectroscopy [3]. A comparison of the oxidation potentials of aromatic compounds with alcohols leads to the same conclusion. For instance, the half-wave potential E1/2 of toluene is 1.98 V vs. Ag/Ag+, whereas for

TAbLE 26.1 Half-Wave Oxidation Potentials E1/2 of Several Aliphatic, benzylic, and Allylic Alcohols Obtained by Cyclic Voltammetry

Entry

Compound

E1/2 vs. Ag/0.01 M Ag+ (V)

References

Entry

a

[2]

5

1

MeOH

2.69

2

EtOH

2.57a

[2]

6

3

iPrOH

2.46a

[2]

7

[5]

8

E1/2 vs. Ag/0.01 M Ag+ (V)

Compound OH

OH

References

1.59

b

[5]

1.22/1.64b

[5]

1.31b

[5]

O

4

OH

>2.0b

Solvent, acetonitrile; working electrode, platinum. Supporting electrolyte: a 0.15 M NBu BF . 4 4 b 0.5 M NaClO . 4

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OH

OH

>2.0b

[5]

983

Oxygen-Containing Compounds CPE divided cell CH3CN/LiClO4

OH

R = o - OMe, p-OMe: 60% R = m-OMe: iPr > * The abbreviations CPE and CCE will be used for controlled potential and controlled current electrolysis throughout this chapter.

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984

Organic Electrochemistry

O

CPE divided cell CH3CN/NBu4BF4 K2CO3

OH R

CPE divided cell CH3CN/NBu4BF4 K2CO3

R R = Me: 83% R = Et: 75%

R = iPr: 63% R = tBu: 61%

SCHEME 26.3

O

Preparative scale electro-oxidation of different secondary benzyl alcohols.

Et > Me. Tertiary alkylphenyl alcohols were found to cleave alkyl fragments to yield the respective ketone [7]. The nature of the leaving group for path 2 was studied in more detail. The liberated carbocation R+ could undergo a Ritter-type reaction with acetonitrile to form the corresponding amide. For radical R•, either H abstraction from the solvent or further oxidation with subsequent Ritter-type reaction is conceivable. When 1-methyl- and 1-ethyl-substituted benzyl alcohols were electrolyzed, the formation of methane and ethane was observed exclusively. This indicates that the course of the reaction is more likely represented by path 2a. However, after the conversion of iPr- and tBu-substituted benzyl alcohols, the corresponding amides can be found as by-products. As iPr- and tBu-radicals have a much longer lifetime, they can be anodically oxidized to undergo the Ritter-type reaction and the formation of the respective amides was indeed confirmed. The electrochemical oxidation of naturally occurring allylic alcohols is despite their abundance only scarcely studied. For instance, when geraniol or crotyl alcohol is directly electrolyzed in an acetonitrile/methanol mixture containing pyridine and LiClO4 (undivided cell, galvanostatic conditions), the corresponding α,β-unsaturated aldehydes (citral and crotylaldehyde) or, respectively, their dimethyl acetals are obtained in good yields [8]. In contrast, electrolysis of prenol unexpectedly leads to γ-hydroxy acetal 3 in 75% yield and not to prenyl aldehyde (4) (see Scheme 26.4). An unusual rearrangement starting from 4 and presumably occurring at the electrode surface was proposed. The initial oxidation step of the sequence depicted in Scheme 26.4 was confirmed, since subjection of 4 to the same electrolysis conditions equally leads to the formation of 3. The anodic treatment of 1,2-glycols in methanol containing tetraethylammonium p-toluenesulfonate at potentials between 1.7 and 2.3 V vs. SCE leads to the oxidative cleavage under the formation of the carbonyl compounds or acetals (Scheme 26.5) [9]. This method is quite useful compared to methods involving Pb(OAc)4, NaIO4 or O3, since a tedious work-up is avoided and no toxic reagents are employed. Moreover, no stereochemical limitations, which typically occur using conventional methods [10], were found. 1-Hydroxy-2-methoxy and 1,2-dimethoxy compounds can be transformed in similar current efficiencies to those of the respective diols.

OH

CCE undivided cell CH3CN/MeOH/ pyridine 10:10:1 LiClO4

MeO

OH 3 75%

OMe

MeOH MeOH MeO

O O 4 H+

SCHEME 26.4

Anodic oxidation of prenol.

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O

985

Oxygen-Containing Compounds

HO

O

CPE MeOH/p–TsONEt4

R

O

R

R

+

n

O

O

O

HO

+

n

n

R

O

O

R

R

n

O

O

R = H, Me, iPr n = 3, 4 35– 96%

SCHEME 26.5

Anodic conversion of cyclic 1,2-diols. CCE undivided cell NEt4BF4

OH F 3C

F3C

O

CF3

OEt

F3C

F3C

5 >70%

SCHEME 26.6

OH

O

OH

6

7

Electrochemical oxidation of 2,2,2-trifluoroethanol.

R1

O n

OH

CCE undivided cell EtOH/NaOEt/LiBF4

R2

R1

n

O O

R2

R1 = C6H13, (C2H5)2CH R2 = H, CH3

8 51–61%

n = 1, 2

SCHEME 26.7

Oxidative cyclization of ω-hydroxy tetrahydropyranes.

Despite the high redox stability of fluorinated alcohols toward oxidation [11–14], 2,2,2-trifluoroethanol can be converted under galvanostatic and solvent-free conditions to 2,2,2-trifluoroethyl hemiacetal 5 (see Scheme 26.6) [15]. The hemiacetal represents a useful alternative to fluoral (6), which is an important building block for compounds exhibiting trifluoromethyl groups. However, fluoral is difficult to handle due to its low boiling point [16–18]. In laboratory-scale preparations, the commercially available ethyl hemiacetal 7 is mostly employed as a substitute for 6. It was demonstrated that 5 has a better reactivity compared to 7 when employed for Grignard reactions or acidcatalyzed electrophilic aromatic substitutions [15]. Oxidative cyclization of ω-hydroxy tetrahydropyrans under full stereocontrol can be achieved by the anodic oxidation of the starting material in the presence of a suitable base (see Scheme 26.7) [19]. An electrolyte consisting of NaOEt and LiBF4 in ethanol provides the best conditions for the preparation of the desired [4,5] and [5,5] spiroketals (8), which represent a common structural motif in numerous natural products and pharmaceutics [20]. A plausible mechanism involves initial deprotonation of the ω-hydroxy functionality followed by anodic oxidation to alkoxyl radical 9. Subsequent intramolecular hydrogen abstraction and further oxidation lead to oxonium ion 10, which is attacked by the ω-hydroxy group to form the desired spiroketal 8 (Scheme 26.8).

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986 R1

Organic Electrochemistry O

OH

R1

NaOEt

n

– O

O

n

–e–

O H

R2

R2

R1

n

O

R2 9

R1

O

n

OH

–e–

+ O

R1

n

R2

R1

OH

n

O O

R2

R2

10

8

SCHEME 26.8 Proposed mechanism for the oxidative cyclization of ω-hydroxy tetrahydropyranes.

2. Indirect Oxidation Compared to direct oxidation of alcohols, mediated electrolysis renders by far superior results in terms of selectivity, compatibility with functional groups, and energy efficiency. For instance, the employment of iodide salts under galvanostatic conditions allows for efficient transformation of aliphatic primary and secondary alcohols to the corresponding carboxylic acids and ketones, respectively [21]. It should be pointed out that such transformations are not possible with direct electrolysis. Furthermore, when iodide salts are used as mediators, the addition of base is not necessary. Other successfully tested mediating systems for the oxidation of aliphatic alcohols are alkali metal nitrates in acetonitrile or in biphasic solution [22,23], thioanisole in benzonitrile or 2,2,2-trifluoroethanol [24,25], and a double mediatory system consisting of RuO4/RuO2 and Cl−/Cl+ in a biphasic system [26]. A very successful and convenient approach for the oxidation of all types of alcohols is the employment of Ni(OH)2 electrodes in aqueous alkaline media (see Scheme 26.9) [27–30]. In this type of reaction, the electron transfer to the alcohol substrate is heterogeneously catalyzed by the anode material, which means that no mediator has to be separated from the reaction mixture. The scope of this method comprises numerous oxidations of primary and secondary aliphatic and benzylic alcohols [27,29], oxidative cleavage of vicinal diols, and chemoselective oxidation of hydroxylsteroids and partially protected sugars [27,28,30]. The nickel electrode has to be activated prior to electrolysis by the deposition of a thin Ni(OH)2 film from an aqueous basic Ni(II) salt solution. From this film, a black surface layer of Ni(III) oxide hydroxide is continuously electrogenerated during electrolysis (see Scheme 26.10) [31,32]. After the rate-determining radical hydrogen abstraction, an α-hydroxyl radical is formed that is readily oxidized

RCH2OH + 5OH–

CCE Ni(OH)2 anode

O R C O– + 4H2O + 4e–

H2O/NaOH

SCHEME 26.9 Alcohol oxidation on the Ni(OH)2 electrode.

Ni(OH)2

+ OH–

NiOOH + RCH2OH H R C OH

SCHEME 26.10

0.6 V vs. SCE

NiOOH + H2O + e–

Ni(OH)2 + R

H C

OH

O R C O–

Mechanism for electro-oxidations on the Ni(OH)2 electrode.

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987

Oxygen-Containing Compounds

to the ketone in the case of secondary alcohols and to carboxylic acid in the case of aldehydes [27]. The formation of the catalytically active film already takes place around 0.6 V vs. SCE, which is a dramatic decrease compared to the potentials for direct alcohol oxidation depicted in Table 26.1. Hence, functional groups such as ester, alkyne, or electron-rich aromatic rings such as furans are tolerated. The most frequently studied type of mediatory system for the oxidation of alcohols is the class of N-oxyl radicals. Particularly, 2,2,6,6-tetramethylpiperidinyl-N-oxyl (TEMPO) and corresponding derivatives are well studied. Basically, the employment of N-hydroxyphthalimides renders similar results [33]. In contrast to the use of such mediators in conventional organic synthesis, the employment of a stoichiometric amount of oxidant (e.g., NaClO) is avoided. This indirect method allows for a selective transformation of primary and secondary alcohols to carbonyl compounds at very low potentials (typically around 0.4 V vs. Ag/AgNO3) [34,35]. Overoxidation of aldehydes to carboxylic acids can be avoided in aprotic media [34], whereas electrolysis in aqueous basic solution generally affords the carboxylic acid [35]. The method requires potential control and operation in a divided cell [34]. Slow reaction of secondary alcohols allows for a selective transformation of primary alcohols in the presence of secondary hydroxy groups. Further advantages are rapid conversion at low temperatures of up to –60 °C and high turnover rates of the catalyst. The reaction can be conducted either homogeneously in solution or with immobilized TEMPO on a modified carbon felt electrode [34,36]. Despite the classical organic and aqueous electrolyte systems, further electrolysis media such as oil-in-water nanoemulsions and ionic liquids were found to be suitable [37,38]. By anodic oxidation of the nitrosyl radical, the active oxoammonium ion is formed, which reacts with an alcohol to give the corresponding hydroxylamine and carbonyl compound (see Scheme 26.11). Reasonable reaction rates are only achieved in the presence of a suitable base such as lutidine. In aqueous systems, a buffer system such as K2CO3/KHCO3 is useful. Direct anodic regeneration of the active species requires a potential of 0.8 V vs. Ag/Ag+. As the hydroxylamine undergoes comproportionation with the oxoammonium compound to give the N-oxyl radical, the catalytic cycle can be set up at the lower oxidation potential of the latter species. A significant improvement of the method described earlier is the employment of a double mediatory system in biphasic media (see Scheme 26.12), wherein the active bromine species is generated anodically in the aqueous phase to react with the N-oxyl radical to form the oxoammonium species [39]. The conversion of alcohols to carbonyl compounds then proceeds in the organic phase. The advantage of this method is a simple experimental setup. In contrast to the single mediatory system (see Scheme 26.11), the operation in an undivided cell at galvanostatic conditions within a wide range of current densities is possible. Based on the initial studies carried out with TEMPO, numerous modifications of this type of mediator were designed in order to tailor the system for specific purposes (see Figure 26.1). From the environmental point of view, the employment of water-soluble TEMPO derivatives 11a and 11b (WS-TEMPOs) R 2

N

R

0.4 V vs. Ag/AgNO3

2

O–

R + R N

O

O

R΄

0.8 V vs. Ag/AgNO3

Comproportionation

R + O N R O H R

N

R΄

R΄

R

OH + R + R N

O R΄

O

SCHEME 26.11

Electro-oxidation of alcohols mediated by N-oxyl radical species.

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OH

Base

988

Organic Electrochemistry

Anode

Br– + 2OH–

R 2 +N O R

Alcohol Carbonyl compound R N R

2e–

OBr– + H2O

OH

R 2 N O R

Aqueous Organic phase phase

SCHEME 26.12

Double mediatory system in biphasic medium based on TEMPO and bromide salt. H N

Y

iPr N

O

O –

11a:

Y = –NR3+ Br

11b:

Y = –SO3– H

+

FIgURE 26.1

N

O N

O

N O

R

AcHN

Cl 12

14

13

Several developments in the field of N-oxyl-based mediators for electrooxidation of alcohols.

is attractive, since they allow for the use of aqueous electrolytes [40,41]. With this type of mediator, the scope is not just limited to water-soluble alcohols, as nanoemulsions are formed upon ultrasonic irradiation of a solution of WS-TEMPO and lipophilic alcohol in aqueous electrolyte. After electrolysis, the resulting aldehyde can easily be isolated by extraction with an organic solvent. The aqueous electrolyte can be reused for several cycles, since the concentration of the mediator remains almost unaffected. When an optically active N-oxyl species is used for the oxidation of racemic sec-alcohols, a kinetic resolution can be accomplished. In this case, the product mixture contains ketone and enantiomerically enriched alcohol. This concept was first developed for the conventional oxidation method and later transferred to the electrocatalytic reaction using 12 as optically active mediator [42–44]. The employment of a double-mediatory system such as sodium bromide combined with 13 in biphasic solution again proved to be beneficial (see Scheme 26.13), allowing for operation in an undivided cell under galvanostatic conditions [45]. Similar results can be obtained in an organic solvent-free method using an aqueous silica gel supported system [46]. The enantioselective oxidation of racemic sec-alcohols mentioned earlier can also be achieved upon the employment of an optically active amine base such as (−)-sparteine in combination with a TEMPO-modified graphite felt electrode [47,48]. The slow reaction with sec-alcohols as a typical feature of TEMPO derivatives is clearly advantageous for the discrimination between hydroxy groups at different positions of a molecule. However, it turns into a drawback when the oxidation of sterically hindered alcohols such as (−)-menthol is intended (see Scheme 26.14) [49]. OH R R = aryl

SCHEME 26.13

CCE (3 F ) undivided cell NaBr/cat. 13 CH2Cl2 aq. NaHCO3

O R 42–61%

OH +

R 23–55% ee: 54–91%

Kinetic resolution of racemic alcohols using optically active N-oxyl radical species 13.

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989

Oxygen-Containing Compounds

CCE 14 or TEMPO CH2Cl2 sat. aq. Na2CO3

OH

O

TEMPO: 23% 14: 99%

SCHEME 26.14 R R

Indirect electro-oxidation of (─)-menthol using N-oxyl radical species 14 or TEMPO. O

Br2

OH

R

O SnBu2

Bu2SnO

O

R

15

R

OH

CCE undivided cell

R

O

R

OH

cat. Bu2SnCl2 Et4NBr/MeOH

R

OH

R = alkyl up to 96%

Conventional method

SCHEME 26.15

Selective anodic oxidation of vicinal diols to α-hydroxy ketones.

Azabicyclo-N-oxyls 14 offer the advantage of a well-accessible N-oxyl group and exhibit a superior performance compared to TEMPO when sterically demanding groups adjacent to the secondary hydroxy functionality are involved. Analogous to 12 and 13, optically active versions of 14 are efficient electrocatalysts for kinetic resolution of racemic sec-alcohols (compare Scheme 26.13) [50]. Common oxidizing reagents or methods such as Pb(OAc)4, IO4−, Swern conditions, or direct electrolysis lead to the oxidative cleavage of vicinal diols or oxidation to the diketones, respectively [9,51–54]. Hence, the selective conversion of 1,2-diols to α-ketoalcohols is a quite challenging task. Nonelectrochemically, it can be accomplished by conventional methods when a vicinal diol is transformed to the corresponding stannylene acetal (15) followed by subjection to brominolysis (see Scheme 26.15, left) [55]. However, this method suffers from some disadvantages such as the employment of excess Bu2SnO and addition of bromine, which often causes side reactions and requires a tedious work-up. An elegant way to circumvent these restrictions is the indirect anodic oxidation using alkylammonium bromides as mediator/supporting electrolyte and catalytic amounts of Bu2SnCl2 (see Scheme 26.15, right) [56]. In a plausible mechanism (see Scheme 26.16), the diol reacts with (OMe)2SnBu2 to form stannylene acetal 15 or the related zwitterion 16. After reaction with the anodically formed oxidant (Br+) R

OH

R

O

Cl2SnR2

(MeO)2SnR2

Br–

–2e– anode

R

+ OSnR2

R

OBr

2MeO– –1/2H2

+ 2e– cathode

R

OH

R

OH

2 MeOH

+ Br R

+ OSnR2

R

O

R

O–

R

O

SnR2 16

SCHEME 26.16

15

Proposed mechanism for the anodic oxidation of vicinal diols.

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990

Organic Electrochemistry

and cathodically formed base (MeO −), the hydroxyketone is obtained and the organotin catalyst will be regenerated. The method was successfully applied to a broad scope of 1,2-glycol derivatives, and some important characteristics of this type of reaction were concluded [57]. In terms of selectivity, it was found that primary and tertiary alcohols cannot be oxidized using this method. This feature can be exploited for the determination of the regioselectivity, when primary/secondary and tertiary/ secondary alcohols are oxidized. Furthermore, in the presence of a third hydroxy group as, for example, in glycerol, a high selectivity for 1,2-diols was observed. In the case of two secondary hydroxy groups, the sterically less hindered one was predominantly oxidized.

B. REDUCTION The transformation of alcohols into the corresponding alkanes by the removal of the hydroxy function is of high importance in organic chemistry, and selective conversion in the presence of sensitive functional groups still remains a challenge. Aside from other multistep sequences [58–60], the Barton–McCombie reaction and its variants are most frequently employed for deoxygenation on laboratory scale [61,62]. However, this type of reaction involves some serious limitations such as the employment of toxic organotin reagents and the preparation of the sensitive xanthates. Similar to direct anodic transformations, the cathodic conversion of alcohols to alkanes is difficult, as very negative potentials have to be applied. For catalytically active electrode materials, proton discharge under the formation of the hardly reducible alcoholate is the preferred reaction. In contrast, the employment of materials with high overpotential for hydrogen evolution such as mercury, lead, or carbon allows for a direct deoxygenation in very few examples [63,64]. At the required potentials of up to –2.9 V, the presence of most functional groups is not tolerated (see Table 26.2). Furthermore, only benzylic and allylic alcohols are cathodically active within the potential windows of common organic electrolytes. The scope of possible substrates for direct cathodic deoxygenation is therefore narrow. However, the oxidation potentials can be significantly decreased when the oxygen is functionalized into suitable leaving groups.

TAbLE 26.2 Half-Wave Reduction Potentials E1/2 of Several benzylic, Allylic, and Propargyl Alcohols Obtained by Polarography Entry

Compound

1

E1/2 vs. SCE (V)

Entry

OH

—

4

OH

–2.90

5

–2.81

6

Compound

E1/2 vs. SCE (V) –2.44

OH

Ph

2

–2.57/–2.80

Ph

Ph 3

OH

Ph

Ph Ph

OH

OH Ph

Ph Source: Lund, H. et al., Electrochim. Acta, 19, 629, 1974. Solvent, DMF; supporting electrolyte, 0.1 M NBu4I. Working electrode: mercury.

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–2.74

991

Oxygen-Containing Compounds

1. Direct Reduction The cathodic deoxygenation proceeds selectively on mercury or lead electrodes when the hydroxy functionality is situated in benzylic position [63]. As the required potentials are very high, further C–C double or triple bonds in the substrate are reduced as well (see Scheme 26.17). As most other functional groups are not tolerated at these harsh conditions, the practical use of this method is rather limited [65]. 2. Cathodic Deoxygenation after Introduction of Leaving groups A more promising approach for C–O bond cleavage is the employment of leaving groups that significantly lower the reduction potential of the substrate (see Scheme 26.18). Typically, these leaving groups are anions of strong acids. For instance, methanesulfonates can be converted under galvanostatic conditions using a lead cathode. The transformation into the corresponding deoxygenated product proceeds selectively and tolerates several functional groups such as ester, nitrile, and epoxide [66]. In the case of 1,3-diols, the electroreduction proceeds with the formation of cyclopropanes according to Scheme 26.19 [67]. As oxalic acid esters are easily reduced (around –1.4 V vs. Ag/AgI), their leaving group ability was exploited for the deoxygenation of a broad range of benzyl alcohols (see Scheme 26.20) [68,69]. No additional preparative step is required, since the active species is formed in the electrochemical cell via base-catalyzed transesterification. The course of this reductive cleavage was studied in more detail with coulometry, cyclic voltammetry, and product analysis, and the mechanism depicted in Schemes 26.21 through 26.25 was proposed [69]. As initiation step, the electrogenerated base 17 deprotonates the substrate ROH that subsequently undergoes transesterification (Scheme 26.22). Ester 18 is reduced in the next step followed by decomposition into benzyl radical (19) and leaving group EtOCOCO2− (Scheme 26.24). The fate of the radical is most probably proton abstraction from a solvent molecule S–H (Scheme 26.25), since the consumption of only one F was observed and further reduction and protonation would require a second F. As radical 19 needs to be stabilized by conjugation, the scope of this method is restricted to benzyl alcohols. OH

CPE mercury cathode

R1

SCHEME 26.17

R2

DMF/NBu4I

R2

R1 = R2 = Ph: 95% R1 = Ph, R2 = Me: 90% R1 = R2 = Me: 0%

R1

Electrochemical reduction of alkynols. OX R R΄ X = leaving group

OH

H H

R΄

R

R

R΄

Difficult

SCHEME 26.18

Electro-reductive deoxygenation of alcohols. R1 OMs R2 OMs

CCE divided cell DMF/TsONEt4

R3 R2 = alkyl, H R1, R3 = alkyl, alkoxy

SCHEME 26.19

R1 R2 R3 Up to 97%

Electrochemical synthesis of substituted cyclopropanes.

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992

Organic Electrochemistry CPE divided cell DMF/Bu4NClO4

OH R1

O

R2

EtO

R1 = aryl R2 = aryl, alkyl

R1

OEt

R2

O Up to 97%

SCHEME 26.20 Electro-reductive deoxygenation of alcohols using oxalic acid ester.

Initiation: EtO

O

R

OH

EtO

OH

O

OEt

+ OEt

O

OEt

O

O–

EtO

+e–

R

O–

17

SCHEME 26.21

Initiation step for the deoxygenation of alcohols using oxalic acid ester.

Base-catalyzed transesterification: EtO

O +

R

O

O

O

R

OET

O

EtO O–

+ EtO–

18

SCHEME 26.22 Base-catalyzed transesterification. EtO–

+ R

EtOH

OH

+ R

O–

SCHEME 26.23 Deprotonation.

Cathodic cleavage: EtO

O R

EtO

+e–

EtO

O

R

O

O

O–

O

+ R – CH2

O

O

18

O– 19

SCHEME 26.24 Cathodic cleavage.

Termination: 19

+

S H

R CH3

+ S

SCHEME 26.25 Termination step.

The employment of toluates or diphenylphosphinates as leaving groups represents a significant improvement of this type of reaction (see Scheme 26.26) [70–72]. With both types of esters, the scope of the reaction is broadened compared to the oxalate-based method, since the hydroxy functionality of the substrate must not necessarily be situated in benzylic position. Similar to the process depicted in Schemes 26.24 and 26.25, the cathodic reduction of such esters leads to the formation of a radical cation followed by the decomposition into alkyl fragment and anion. The desired product is then formed by proton abstraction from the solvent. As only 1 F is consumed and the nature of the leaving group can unambiguously be identified, an EC type mechanism

© 2016 by Taylor & Francis Group, LLC

993

Oxygen-Containing Compounds

R O R΄

CCE divided cell DMP or DMF NBu4BF4

–

R = alkyl, benzyl;

O–

R H + R΄

R O R΄

R΄ = C(O)aryl, P(O)(Ph)2

SCHEME 26.26 Electrochemical reduction of toluates and diphenylphosphinates.

R OH + PR΄3 R = alkyl, benzyl R΄ = Ph, Bu, OPh

CCE undivided cell CH3CN/Et4NBr

R΄ R H

+ O=PR΄3

via R

48–96%

+

O

R΄

R΄ 20

SCHEME 26.27 Electrochemical deoxygenation of alcohols in the presence of tertiary phosphine.

can be proposed. For the studied reactions, the decomposition rate can be clearly correlated with the stability of the produced radical [71]. Consequently, diphenylphosphinate radical anions exhibit a much faster decomposition rate compared to the toluate species. This is advantageous from the preparative point of view, since this highly reactive species is less prone to side reactions. For secondary and tertiary alcohols, the toluate ester has to be prepared separately prior to electrolysis [70], whereas in situ transesterification in the presence of methyl toluate is the method of choice for primary alcohols [72]. The diphenylphosphinates are prepared ex situ for all kinds of alcohols [71]. Primary and secondary alcohols can also be efficiently deoxygenated by a double electrolysis in the presence of phosphine (see Scheme 26.27) [73,74]. The active phosphonium species 20 is formed in situ by anodic oxidation of the phosphine and subsequent reaction with the alcohol. 20 is then cathodically reduced to give the alkane and phosphine oxide. The scope of this reaction comprises benzylic and aliphatic alcohols. In contrast, tertiary alcohols cannot be deoxygenated at these conditions. As this method requires no basic conditions, ester groups are tolerated. Although some of the electrolytic conversions seem to be limited for aliphatic alcohols, there are two major developments that might overcome the current limitations: Novel electrode materials such as boron-doped diamond (BDD) allow the direct generation of oxyl radical by anodic treatment or the cathodic operation at highly negative potential [75,76]. For nondestructive purposes, these pathways are yet not well investigated. Another strategy will be the development of novel and highly specific mediators to circumvent a direct conversion at the electrode surface [77].

II.

ELECTROCHEMICAL CONVERSION OF ETHERS

The ether moiety is a very common motif and therefore omnipresent. The pronounced electrochemical stability of saturated aliphatic ethers and their ability to dissolve salts allow the usage of ethers as a solvent component in the electrolyte. Therefore, unsaturated ethers are much more easily converted. However, ethers are in particular prone to hydrogen atom abstraction, since the intermediate radical experiences stabilization by the oxygen atom.

A. OXIDATION Similar to aliphatic alcohols, the direct anodic oxidation of saturated monoethers proceeds at relatively high potentials. For instance, THF is electrochemically inert up to 2.0 V vs. SCE [78]. Hence, only few examples for preparative anodic conversions of saturated ethers can be found in literature as many functional groups are not tolerated at such high potentials [79,80].

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994

Organic Electrochemistry

TAbLE 26.3 Half-Wave Oxidation Potentials E1/2 of Several Aliphatic, Phenyl, and Vinyl Ethers Obtained by Cyclic Voltammetry Entry 1

Compound

E1/2 vs. SCE (V) >2.0

a

O

References

Entry

[78]

4

Compound MeO

2

OMe

1.76b

[85]

5

OMe

1.45b

[85]

6

References

1.42

b

[85]

1.12b

[85]

1.4c

[86]

OMe OMe

OMe

MeO 3

E1/2 vs. SCE (V)

O

OMe Solvent, acetonitrile; working electrode, platinum. Supporting electrolyte: a Electrolyte, THF/0.65 M NEt BF . 4 4 b Electrolyte, CH CN/0.1 M Pr NClO . 3 4 4 c Value refers to peak potential; Electrolyte, CH NO /1 M LiClO . 3 2 4

In contrast, unsaturated ethers cover a wide range of oxidation potentials and are generally easier to oxidize compared to their saturated congeners (see Table 26.3, entries 2–6). Consequently, the preparative electrochemistry of such compounds is well explored and their reactivity exploited for numerous useful synthetic transformations [81–84]. 1. Saturated Ethers Anodic oxidation of aliphatic ethers typically leads to oxonium ion 21, which is subsequently attacked to give the α-substituted product (see Scheme 26.28). For instance, electrolysis of nonsubstituted, cyclic, and acyclic ethers in basic methanolic or aqueous solution allows for the oxidation to the corresponding acetals and hemiacetals in moderate yields [87,88]. Significantly better results for the oxidation of ethers to acetals can be obtained when a solvent mixture of methanol and acetic acid is used (see Scheme 26.29) [79]. An intriguing example for such reactivity was reported more recently in the context of the threecomponent synthesis of protected homoallylic alcohols in THF-based electrolyte (see Scheme 26.30) [80].

R1

O

R2

–2e–, –H+ R1

O

R2

Nu

Nu R1

O

R2

21

SCHEME 26.28 Anodic oxidation of aliphatic ethers. CCE undivided cell CH3OH/AcOH

O n

76% (n = 1) 71% (n = 2)

O

OMe n

SCHEME 26.29 Preparative scale electro-oxidation of cyclic ethers to the corresponding acetals.

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995

Oxygen-Containing Compounds

O Ph

+

CCE undivided cell

O

Br +

O

O Ph

Indium cathode platinum anode

22 60–70%

SCHEME 26.30

Electrochemical three-component synthesis of protected homoallylic alcohols. O

ln III

Anode

Cathode Br

+ O

Oln III ln0 Ph Ph

O

23

24

O

O Ph 22

SCHEME 26.31

Proposed mechanism for the three-component synthesis of homoallylic alcohols.

In this process, benzaldehyde, allylbromide, and THF react under indium mediation in a convergent paired electrolysis at galvanostatic conditions to form 22 in good yields. The course of the reaction is depicted in Scheme 26.31. Reaction of allyl bromide with the indium metal of the cathode affords an allylindium species that reacts with benzaldehyde to form homoallyl alcoholate 23. During electrolysis, the In(III) salts generated by the chemical reaction are constantly reduced, while the solvent THF is anodically oxidized to oxonium ion 24, which reacts with 23 to give homoallyl acetal 22. The application of indium metal as cathode material is crucial for the process, since the reduction of In(III) salts on typical electrodes such as platinum or glassy carbon proceeds at very negative potentials. Hence, the employment of catalytic amounts of In(I) salts in combination with a standard electrode material is not possible. In order to favor the oxidation of the solvent, high current densities at the anode have to be realized by using an undersized platinum wire (quasidivided cell) [89]. 2. Unsaturated Ethers In contrast to saturated ethers, the oxidation of vinyl ethers proceeds at relatively low potentials (compare Table 26.3). Typically, this reaction proceeds in a single-electron transfer step, leading to the relatively stable radical cation 26 according to Scheme 26.32 [81]. The type of the subsequent reaction strongly depends on the properties of the electrolyte. In strongly basic or nucleophilic Path a

RO

RO

Nu–

Nu

Nu

Nu

RO –e–

RO

RO +

+

RO

Path b

OR

Nu

Nu– 25

26

RO RO Path c

+

RO +

SCHEME 26.32

OR

Typical pathways for the electro-oxidation of vinyl ethers.

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OR

+ Nu

996

Organic Electrochemistry CCE undivided cell OR R = alkyl

SCHEME 26.33

OMe MeO

MeOH/KOH –5°C

OR 27 40–75%

Preparative scale electro-oxidation of vinyl alkyl ethers.

environment, the electrophilic reaction of 26 is preferred (path a), whereas under neutral or slightly basic conditions, the fate of 26 is determined by radical dimerization (path b) or electrophilic addition to 25 (path c), both leading to dimers of 25. For instance, anodic treatment of aliphatic vinyl ethers in methanolic KOH renders predominantly the methoxylated product 27 (see Scheme 26.33) [90]. Under these conditions, the yield of the desired product is significantly impaired by the dimerization reaction (yield of the dimer, 15–24%). The potential use of this side reaction as a C–C bond formation method was realized soon after, followed by the optimization of the electrolysis conditions for this purpose [81]. In this context, it was found that under less basic conditions, the conversion of aliphatic vinyl ethers leads preferentially to the dimerization product 28 (Scheme 26.34). The use of a MeOH/NaClO4 electrolyte containing 2,6-lutidine proved to be beneficial. The addition of a weak base is essential and serves the trapping of protons, which are released upon nucleophilic attack of methanol on the cationic or radical cationic dimer intermediate. When the reaction is carried out in the absence of base, the concurrent proton-catalyzed acetalization of the enol ether significantly lowers the yield of the desired product. The C–C bond forming reaction discussed earlier was also used for intramolecular cyclizations (Scheme 26.35) [82]. Using similar electrolysis conditions as depicted in Scheme 26.34, five-, six-, and seven-membered rings can be obtained from the corresponding bis(enol ether) substrates in reasonable yield without stereoselectivity. It was found that the method is unsuitable for the synthesis of compounds with larger ring sizes and that the yield decreases with increasing chain length. Moreover, it was demonstrated that this method allows for the generation of one or even two vicinal quaternary carbon centers upon cyclization [82]. Silyl enol ethers proved to be similarly suitable substrates for this type of electroorganic transformation [82,91]. This reaction principle was extended to intramolecular radical-cation cyclizations of vinyl ether substrates containing electron-rich aromatic rings as nucleophilic element (Scheme 26.36) [92]. The product selectivity was observed to be strongly influenced by the position of the activating groups on the aryl moiety. For instance, constant current oxidation of m,m-dimethoxy-substituted 29 leads R1O

R4

R2

R3

2

R1 = Me, Et R2, R3, R4 = H, alkyl

SCHEME 26.34

CPE undivided cell CH3OH/NaClO4 lutidine –10°C

R4

R2 OR1 MeO OMe R1O 4 R2 R3 R 28 Up to 61% R3

Electrochemical dimerization of enol ethers. MeO R R n R = H, Me n = 1, 2, 3

SCHEME 26.35

OMe

CCE undivided cell

OMe

CH3OH/CH3CN NaClO4/lutidine

OMe

R R n MeO 50–70%

Anodic intramolecular cyclization of enol ether substrates.

© 2016 by Taylor & Francis Group, LLC

OMe

997

Oxygen-Containing Compounds MeO R1 = R3 = H R2 = OMe

OMe MeO

30 31%

OMe CCE undivided cell CH3OH/CH2Cl2 LiClO4/lutidine

R2

R1

R1 = OMe R2 = R3 = H OMe

R2 R3

31 51%

O

OMe

OMe R1 = R2 = H R3 = Me

29

32 51% O

SCHEME 26.36

Intramolecular cyclization of enol ether substrates containing electron-rich aromatic rings.

to fused bicyclic structure 30, whereas under the same electrolysis conditions, the conversion of p-methoxy-substituted 29 affords spirodienone 31. In the presence of two or more activating substituents on the ring, overoxidation of the desired products was found to significantly decrease the yield. The resulting limitation of the scope of substrates can be circumvented using vinyl thioether derivatives of 29. When a nonactivated aryl ring is appended, a different reaction path is opened up. In this case, elimination of a proton in α-position to the radical cation leading to α,β-unsaturated ketone 32 preferentially occurs. Hence, the presence of electron-donating groups seems to be crucial for the cyclization reaction. Additionally, it was demonstrated that other electron-rich aromatic compounds such as furan and pyrrole are also suitable for this type of cyclization reaction [92]. The electrochemical conversion of enol ethers in [2+2] cycloaddition reactions was explored more recently [84,86,93,94]. In a plausible mechanism, the anodically generated radical cation attacks an α-aryl-substituted olefin under the formation of cycloadduct intermediate 33 (Scheme 26.37).

+

Rn

OR 33

Rn

R

+ O

Rn

+ 34

OR

35

OR

Cathode

Anode

R

SCHEME 26.37 tion reactions.

O

Rn

Plausible mechanism for the electrochemical conversion of enol ethers in [2+2] cycloaddi-

© 2016 by Taylor & Francis Group, LLC

998

Organic Electrochemistry R2 R3

R2 R1

CPE LiClO4/CH3NO2

+ R4

O R5

R3

R1

R4 R5

Eox = 1.4 V

O

36 Up to 94% d.r.: up to 25:1

Ri = H, Me, OMe, OPh

SCHEME 26.38

Electrochemical conversion of dihydropyran in [2+2] cycloaddition reactions.

Intramolecular electron transfer proceeds to give 34, which is then reduced to give cyclobutane structure 35. The electron-rich aromatic ring is a crucial component for intermediate stabilization of the cyclobutane radical cation (“redox tag”). A very delightful aspect of this reaction is the possibility of 34 to be reduced by the enolether substrate, which means that only catalytic or substoichiometric amounts of charge have to be passed (typically between 0.1 and 0.5 F). In order to avoid side reactions, the electrolysis should be carried out with controlled potential. The scope of enol ether substrates includes enyloxy benzenes [93], aliphatic acyclic [86], and cyclic enol ethers [84], rendering good to excellent yields. On the olefinic side, numerous unactivated, mono, and gemdisubstituted substrates have been successfully tested. It should be noted that the aromatic redox tag may be attached both to the olefinic substrate as depicted in Scheme 26.37 and to the enol ether substrate [93,94]. The influence of the substitution pattern of the aromatic ring on this type of conversion was studied extensively (Scheme 26.38) [84]. A clear correlation between oxidation potential of the olefinic substrate 36, which is determined by the activating groups on the aromatic ring, and the product yield could be established. As the enol ether substrate has to be oxidized first, the oxidation potential of olefin 36 should be higher than 1.4 V vs. SCE for successful conversion. With potentials above 1.8 V, the yield of the desired product decreases significantly, indicating that the intramolecular electron transfer in intermediate 33 (Scheme 26.37) from cyclobutane radical cation to the aromatic redox tag is inhibited. Among many tested substrates, p-methoxy- or p-phenoxy-substituted phenyl rings rendered the best results with isolated yields of up to 94%. Furthermore, the reaction proceeds diastereoselectively when a cyclic enol ether such as 3,4-dihydro-2H-pyran is employed. The electrochemical conversion of phenyl ethers was studied in the context of protecting group chemistry, particularly in the case of p-methoxyphenyl ether (PMP). This structural element is stable under strongly acidic or basic conditions and has therefore become very useful for the protection of alcohols [95,96]. It is classically cleaved off by the employment of stoichiometric amounts of ceric ammonium nitrate (CAN). Due to the strong oxidative conditions, the presence of several functional groups can be problematic due to degradation of the starting material. In some cases, direct anodic cleavage provides a milder and more selective alternative. The deprotection of several alkyl and allylic PMP ethers was studied with both the electrochemical method and the classic method using CAN as oxidant (Scheme 26.39) [97]. When the electrolysis was carried out in a water OMe

1.7 V vs. SCE NaClO4/NaHCO3 CH3CN/H2O

R

O

R = alkyl, allyl

SCHEME 26.39

Anodic cleavage of p-methoxyphenyl ethers.

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R OH

Up to 92%

999

Oxygen-Containing Compounds

acetonitrile mixture containing NaHCO3 and LiClO4 at 1.7 V vs. SCE using a divided cell, it was found to be superior to the classic method. The orthogonality of this method could be demonstrated, since several acid-, base-, and even oxidation-sensitive protecting groups are tolerated.

B. REDUCTION Generally, ethers are very stable toward electrochemical reduction, and thus, preparatively useful cathodic conversions are hard to achieve. This point becomes clearer in the context of lithium ion battery electrolytes, whereby ethers were initially proposed as solvents [98]. Aliphatic ethers such as THF and diethyl ether are stable on glassy carbon electrodes at 0 V vs. Li/Li+ (–3.2 V vs. Ag/Ag+), whereas despite some further drawbacks, their sensitivity toward anodic oxidation restricts their practical use in energy storage electrolytes [98]. Consequently, only few examples for preparative cathodic reduction of ether compounds are reported in literature. For instance, the cleavage of cinnamyl ethers from conduritol derivatives 37 in the presence of allyl ethers was achieved using a mercury cathode (Scheme 26.40) [99,100]. In contrast to the nonelectrochemical method using sodium in liquid ammonia, the cleavage proceeds selectively in the presence of allyl ethers. The need for a very negative potential (–2.9 V vs. Ag/Ag+) indicates the challenge of such cathodic conversions. In another example, organic electrochemistry was combined with organometallic catalysis in order to achieve efficient conversion of phenyl allyl ethers under galvanostatic conditions to the corresponding phenols in high yields [101]. As active species serves a cathodically generated Ni(0) species, which is constantly formed from the catalytically employed Ni(II) bipyridine complex. A magnesium rod is used as sacrificial anode. It was found that using this cleavage method, several functional groups such as nitril, ester, and halogen atoms are tolerated. Interestingly, for o-(allyloxy) benzaldehydes (38), this reactivity leads to an intramolecular allylation of the carbonyl group by cleavage and transfer of the allyl moiety (Scheme 26.41) [102]. The variety of reaction pathways, which are feasible with ether moieties, allows on the one hand their use as inert tether and on the other hand their selective conversion. This Janus-type structural motif is consequently very valuable for electroorganic synthesis.

R O O

HO

R

R –2.9 V vs. Ag/AgNO3

O O

O

O

OH

O 37

SCHEME 26.40

R = Ph, H

Reductive treatment of conduritol-derived compounds (37). O

OH

Ni(bipy)3(BF4)2 Mg anode Bu4NBF4/DMF

OH

O 38

SCHEME 26.41

87%

Electrochemically induced intramolecular allylic transfer.

© 2016 by Taylor & Francis Group, LLC

O

HO

Ph 78–83%

O

Na/NH3

1000

III.

Organic Electrochemistry

AROMATIC ALCOHOLS (PHENOLS)

Phenols are very common structural entities and in contrast to the aliphatic congeners, the aromatic alcohols represent unique electrophoric groups [103]. Since phenols are usually electron rich and easily deprotonated, the anodic conversion of this particular substrate class has the higher relevance compared to the aliphatic congener.

A. OXIDATION The oxidation potential of phenols is not only strongly affected by substituents on the aromatic core but rather by the nature of the electrolyte employed. Upon the oxidation of 39, the corresponding radical cation is formed. 40 represents a relatively strong acid and a proton will be spontaneously expelled. The anodic treatment of phenols in basic media is facilitated, since the oxidation potential is significantly lowered. The phenoxyl radical 41 has spin densities at several positions as indicated by the mesomeric structures. As a neutral molecule, 41 is not well solvated by a polar electrolyte. In particular, in protic media like sulfuric acid, it will be further oxidized to phenoxonium species 42. If 39 is equipped with electron-releasing groups, the formation is more likely. A radical recombination of 41 is not very likely, since the applied current densities are far too low. Both phenoxyl radical and phenoxonium represent electrophiles and will attack the substrate 39 or other nucleophiles leading to a bond formation. Since different reactive sites are present in intermediate 42, a variety of by-products can be formed apart from the desired 2,2’-biphenol 43. Among these by-products, 2,4’-biphenol 44 and C,O-coupling product 45 play a dominant role. Consequently, the control of the electroorganic conversion of phenols can be challenging [104] (Scheme 26.42).

+

Anode

R

–H+

R

O

O

OH

OH

R

R

40

39

O R

41 Anode

+ O

39 Anode O

O –H+

R

R

R R

42

OH

39 –H+

+ + R

R

OH 44

43 OH O +

+

R 45

SCHEME 26.42

Possible reaction pathway for phenol oxidation.

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OH

OH

R

1001

Oxygen-Containing Compounds

1. Acetoxylation One of the first anodic acetoxylation reactions was reported in 1952 [105]. The authors describe the anodic transformation of naphthalene 46 in acetic acid and sodium acetate to form 1-acetoxynaphthalene 47. The acetoxylation product 47 was hydrolyzed to provide α-naphthol 48 in 24% yield (Scheme 26.43). Scheme 26.44 shows a plausible mechanism for the acetoxylation of toluene 49 [106]. The first step involves the abstraction of an electron from the aromatic hydrocarbon compound to generate an aromatic radical cation 50. This step is followed by the nucleophilic attack of acetate onto the position of the highest positive charge density. In this case, the attack should occur in para-position 51, but other isomers are also possible. A subsequent oxidation leads to the carbenium species 52. Elimination of a proton for rearomatization accomplishes the product 55. The carbenium species allows the migration of the acetyl group by a 1,2-bridging process to form different isomeric products 55–57. So the final isomeric ratio depends not only on the position of the nucleophilic attack but also on the relative energy of the isomeric cation intermediates 52–54. The half-wave potential of the different monoacetoxy derivates can influence the isomeric ratio, too. Table 26.4 displays the half-wave potentials of anisole and its monoacetoxylation products. The potential of 4-acetoxyanisole is about 0.6 V lower than the potential of the starting material [107]. Investigations on anisole under constant potential conditions yield 27% of a mixture of 2- and OAc

OH H2

NaOAc/HOAc

O/OH–

Pt-anode 46

SCHEME 26.43

47

48

Anodic acetoxylation of naphthalene (46). +

–e–

+AcO–

–e– +

50

49

H OAc 51

H OAc 52 –H+

+

OAc H OAc H

53

54 –H+

–H+

OAc OAc OAc 55

56

57

SCHEME 26.44 Plausible mechanism for the acetoxylation of toluene (49).

TAbLE 26.4 Half-Wave Potentials of Anisole and Its Acetoxylation Products Entry 1 2 3 4

Compound

E1/2 vs. SCE (V)

Anisole 2-Acetoxyanisole 3-Acetoxyanisole 4-Acetoxyanisole

1.67 1.74 1.25 1.12

Source: Eberson, L. and Nyberg, K., J. Am. Chem. Soc., 88, 1686, 1966.

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1002

Organic Electrochemistry

CH3

CH3 –e–

49

SCHEME 26.45

+

CH2OAc +AcO–

–e–

–H+ 58

+ CH2

CH2

59

60

61

Mechanism of side-chain acetoxylation of toluene (49).

4-acetoxyanisole with an isomer ratio of 9:1 [107]. Electrolysis with only 50% of the required current gave 40% of the acetoxylation product with an ortho/para ratio of 6:1. This reveals that the para-compound with lower oxidation potential would be consumed at the anode at a higher rate than the other compounds. This directly influences the isomer distribution. Besides the reaction pathway described in Scheme 26.44, loss of a proton can take place at the benzylic position (Scheme 26.45). This step is irreversible and leads to the benzylic radical 59. Another one-electron transfer generates the cation 60, which is attacked by the acetate to form the side-chain acetoxylation product 61. The nuclear isomer ratio for the toluene of 43.2% ortho (57), 11.1% meta (56), and 45.7% para (55) acetoxylation occurs with a ratio of nuclear to side-chain acetoxylation of 71.4–28.6%. The amount of side-chain acetoxylation species depends on different facts: loss of the benzylic proton leads to a benzylic radical with an electron in a p-orbital, which is able to overlap with the aromatic π-system. This elimination is energetically disfavored if the benzylic orbital is sp3-hybridized like the benzylic position in triptycene. In this case, the radical and carbocation species would be highly strained [108]. For this reason, triptycene shows 31.0% acetoxylation in the position 1, 68.5% in the position 2, and only 0.5% in the α-position [109]. One other factor that influences the ratio between side-chain and nuclear acetoxylations is the electrode material. Mesitylene 62 is a good model compound for such studies, because there is only one possible monoacetoxylation product 63 at the aromatic core and one side-chain acetoxylation product 64 (Scheme 26.46). Table 26.5 shows the effect of different anode materials on the ratio between side-chain acetoxylation and the acetoxylation of the aromatic core. The isolated yields are given in Table 26.5 [110]. CH3

CH3 HOAc/NaOAc Anode

62

SCHEME 26.46

CH2OAc OAc +

63

64

Side-chain and acetoxylation of the aromatic core of mesitylene (60).

TAbLE 26.5 Influence of the Electrode Material on the Ratio between Side-Chain Acetoxylation Products and Substitution at the Aromatic Core Entry 1 2 3 4 5 6

Anode Material

Ratio 63/64

Isolated Product (%)

Gold Platinum Rhodium Carbon cloth Glassy carbon Graphite

3.6 4.4 7.6 23 21 23

23 19 57 34 35 56

Source: Eberson, L. et al., Acta Chem. Scand., 29b, 168, 1975.

© 2016 by Taylor & Francis Group, LLC

1003

Oxygen-Containing Compounds

The supporting electrolyte could also influence the product distribution. Acetoxylation of toluene in acetic acid with tetrabutylammonium tetrafluoroborate instead of sodium acetate as supporting electrolyte produces benzyl acetate as major product [110]. Under these conditions, the loss of the benzylic proton of the cationic species seems to be much faster than the nucleophilic attack of acetic acid. Therefore, the substitution on the aromatic core is disfavored. One way to increase the yield of nuclear substitution product is the simultaneous reduction of the α-acetoxylation by-products [111]. This reduction can be performed by adding palladium on charcoal to the electrolyte system. The cathodically generated hydrogen can be consumed to regenerate the starting material. Experiments with p-xylene without Pd/C lead to a ratio of 40% nuclear acetate and 60% α-acetoxylation. Adding only 10% of the catalyst raises the ratio of 2,5-dimethylphenyl acetate up to 98% by the simultaneous increase of total yield from 30% up to 75%. Other examples of acetoxylation are shown at the end of the Section III.A.2. 2. Quinone Formation The anodic oxidation of phenols is frequently studied but represents rather complex chemistry [112]. The different reaction pathways heavily depend on the applied conditions, for example, basic or acidic media, electrolyte, potential, and electrode materials [113]. Sequences of electron and proton transfers lead to different highly reactive intermediates. A well-studied example is the oxidation of 2,4,6-tri-tert-butylphenol (65) [114]. When performing the electrolysis in the presence of a base, the formation of the deprotonated species, the phenoxide ion 66, will happen instantaneously (see Scheme 26.47). After deprotonation, a single-electron transfer leads to the conversion to the phenoxyl radical 67. Under acidic conditions, the removal of one electron generates a radical cation 65a. This radical cation is a rather strong acid. Therefore, a proton transfer occurs and the phenoxyl radical 67 is formed which may further be oxidized to phenoxenium cation 68. A common product of such oxidative phenol treatment is quinone. Phenols equipped with hydroxy groups, para or ortho, yield the corresponding-para or ortho-quinones 69, 70 upon twoelectron transfer (see Scheme 26.48) [115]. Cyclic voltammetry studies were conducted to elucidate the mechanism for the anodic quinone formation. A possible reaction pathway for the quinone generation is depicted in Scheme 26.49. This transformation starts with an initial electron transfer at oxidation potential E1 followed by fast deprotonation to form the radical 72 (see Scheme 26.49). In the second electron transfer, the potential E2 is lower than E1; thus, the oxidation of this radical 72 is easier than that of the hydroquinone 71. The desired quinone 73 is accomplished via another rapid proton elimination [115]. OH

OH t-Bu

t-Bu

+ t-Bu

t-Bu –e– +e– t-Bu

t-Bu 65

65a

–H+

+H+

O–

O t-Bu

t-Bu

–H+

+H+

O t-Bu

t-Bu –e–

t-Bu

t-Bu –e–

+e– t-Bu 66

SCHEME 26.47

t-Bu 67

Anodic oxidation of 2,4,6-tri-tert-butylphenol (65).

© 2016 by Taylor & Francis Group, LLC

t-Bu 68

1004

Organic Electrochemistry OH

SCHEME 26.48

O

Z or

(Y = OH, Z = H, R or Y = H, R, Z = OH) R = alkyl

Y

O

O

–2H+ –2e–

Z

O

Y

69

70

Generation of ortho- and para-quinone. OH

OH –e–

O

O –e–

+

+e– OH 71

O

+e–

–H+ OH

OH 72

E1

–H+

+ OH

O 73

E2

SCHEME 26.49 Proposed mechanism for the quinone formation. OH CN

Cl

CN

Cl OH 74

SCHEME 26.50

O

–2e– –2H+ graphite electrode MeCN/LiClO4

CN

Cl Cl

CN O 75

Reoxidation of 2,3-dichloro-5,6-dicyanohydroquinone (74).

2,3-Dichloro-5,6-dicyano-1,4-benzoquinone (75) is a well-known and synthetically useful oxidant, for example, for the oxidative treatment of alcohols, ethers, or arenes [116]. For synthetic application on larger-scale significant amounts are necessary and problematic waste is produced. From the ecological as well as economical point of view, the reoxidation of 2,3-dichloro-5,6-dicyanohydroquinone (74) to 75 is of interest. The recycling of 74 can be achieved by an electrochemical oxidation in acetonitrile with lithium perchlorate as supporting electrolyte using a divided cell design [117]. Graphite rods as electrode material allow isolated yields of up to 77% (see Scheme 26.50). The anodic conversion of hydroquinones into the corresponding quinones can be interesting in many ways. The generated quinones can serve as intermediates in further transformations like insertions or cycloadditions reactions [118,119]. By this way, complex molecules can be made using theses quinones. A typical application of this hydroquinone conversion is the generation of 3-hydroxy-substituted quinones. After initial electron transfers in aqueous solution from para-hydroquinone 76 equipped with electron-withdrawing moieties like nitro or a carbonyl group, the corresponding para-quinones 77 are provided (see Scheme 26.51). Subsequently, addition of water followed by aromatization generates a trihydroxybenzene 78 [118]. Further anodic two-electron oxidation yields the expected O

OH –2e– –2H+ (R = –CHO, –COCH3, –NO2)

R

OH 76

SCHEME 26.51

OH R

O R

H 2O

–2e– –2H+

R

OH O 77

OH 78

OH O 79

Electrochemical generation of para-quinone with electron-withdrawing groups.

© 2016 by Taylor & Francis Group, LLC

1005

Oxygen-Containing Compounds OH OH

OH

O

OH + O

O

OH

–4e– H2O, NaOAc O

O 80

SCHEME 26.52

81

82

Intermolecular coupling of electrochemically generated ortho-quinone with nucleophiles. OH

O

OH

O

OH

O

OH

OH

OH

O O

O O

O

OH

O

O

O

OH

O

OH OH

O HO

SCHEME 26.53

Coupling products of catechol with β- or α-dicarbonyl compounds.

3-hydroxy-substituted quinone 79. If the starting material exhibits Cl, CH3, OH, or H as substituents in position 2, no insertion of water occurs. Other applications exploit the conversion of ortho-hydroquinone into ortho-quinones. These are highly reactive chemical compounds that can be used as intermediates for intermolecular coupling with different kinds of nucleophiles. Tabaković et al. reported a potential-controlled oxidation of catechol (80) in aqueous sodium acetate including 4-hydroxycoumarin (81) as nucleophile (see Scheme 26.52) [120]. After addition to the ortho-quinone, the intermediate is once more oxidized to a quinoid system and a formal [3+2]-cycloaddition occurs. The desired product 82 of this sequence could be isolated in 95% yields. Scheme 26.53 displays a collection of products obtained by similar treatment of catechol with β- or α-dicarbonyl compounds, respectively [121]. In situ generated quinones can also be used as substrates for Diels–Alder reaction [119]. Scheme 26.54 describes a more elaborated version of such a transformation. The reaction was carried out in an emulsion of 0.5% acetic acid, including water-soluble hydroquinone and water-insoluble  dienes. +

OH HO

OH O CO2Et

O CO2Et

O CO2Et

Nafion

EtO2C

Hydrophobic area

Anode

SCHEME 26.54

In-situ generated para-quinone at Nafion-coated and hydrophobic anodes.

© 2016 by Taylor & Francis Group, LLC

CO2Et

1006

Organic Electrochemistry

TAbLE 26.6 Anodic Formation of Ortho-Quinones with Subsequent Diels–Alder Reaction Entry

Hydroquinone

Diene

Product

yield (%)

83 83 83 84 84

86 87 88 88 85

89 90 91 92 93

88 94 96 94 Quant.

1 2 3 4 5

OH

OH HO

+

R2

R3

83 R1 = Et 84 R1 = C8H17

SCHEME 26.55

–2H+

R4

CO2R1

R2

–2e–

R3 4RO

2C

R4

88 R1 = Et, R2 = R3 = Me, R4 = H 90 R1 = Et, R2 = R3 = H, R4 = CH2CO2Et 91 R1 = Et, R2 = H, R3 = Me, R4 = CH2CO2Et 92 R1 = C8H17, R2 = H, R3 = Me, R4 = CH2CH = CMe2 93 R1 = C8H17, R2 = R3 = Me, R4 = H

85 R2 = R3 = Me, R4 = H 86 R2 = R3 = H, R4 = CH2CO2Et 87 R2 = H, R3 = Me, R4 = CH2CO2Et 88 R2 = H, R3 = Me, R4 = CH2CH = CMe2

Scope of Diels–Alder reaction.

Sodium dodecyl sulfate was used as emulsifier to generate micelles containing the diene. Particularly advantageous is the fact that sodium dodecyl sulfate also serves as supporting electrolyte. For this reason, no further ingredients are required. When performing on bare glassy carbon, only low yields (0–30%) of the desired product were produced. For this reason, the anode was firstly coated with Nafion and after that the surface sulfonic acid moieties were converted to sulfonamide groups to get a hydrophobic layer. Transformations at such type of modified anodes lead to excellent yields up to 96% (see Table 26.6). The conditions described earlier allow the conversion of different hydroquinone and dienes (see Scheme 26.55). Besides the intermolecular conversion of electrochemically generated ortho-quinones, an intramolecular follow-up reaction can take place. Scheme 26.56 depicts the anodic oxidation of tetrahydropapaveroline (94), a dopamine metabolite. Under controlled potential, the corresponding double quinone 95 is formed [122]. In the absence of strong nucleophiles, this product is able to undergo intramolecular Michael addition as subsequent transformation. Instead of hydroquinones, appropriately alkyl-substituted phenols may serve as starting materials to generate dienones. Here, first oxidation steps generate the phenoxonium ion, which represents a highly reactive electrophile. For this reason, it can react with a broad variety of different O

HO

NH

NH

HO

–4e– –4H+

O

OH

O

OH 94

SCHEME 26.56

Electrochemical generation of Michael acceptor.

© 2016 by Taylor & Francis Group, LLC

O 95

1007

Oxygen-Containing Compounds OH R΄

O

O

–2e– –2H+

R˝

R˝

R΄

R΄ or

R˝ Nu

NuH R Nu

R

R

NuH = H2O, ROH, RCO2H, RCN, ArOH, ArOR, ROCH CHR, RSCH CHR

SCHEME 26.57

Mechanism of phenol-oxidation with subsequent nucleophilic attack of impurities.

TAbLE 26.7 Anodic Formation of Cyclohexandienones via Nucleophile Trapping of Phenoxonium Ions Entry Starting Material, Phenol 1 2 3 4 5

2,6-Di-tert-bu-4-Me 2,6-Di-tert-bu-4-Me 4-Methoxy 4-Phenyl 2,6-Di-tert-Bu-4-Ph

Product, Dienone

yield (%)

Anode

Electrolyte System

References

4-Methoxy-4-methyl 4-Methoxy-4-anisyl 4,4-Dimethoxy 4-Methoxy-4-phenyl 4-Ph-4-(l-ala-O-)

95 81 97 87 94

C Pt Pt Pt Pt

MeOH, LiClO4, NaHCO3 MeCN–anisole, Bu4NBF4 MeOH, LiClO4 MeOH, LiClO4 CH2Cl2, Bu4NBF4, 2,6-lutidine

[123] [124] [125] [123] [126]

t-Bu

O

O

OH t-Bu

–2e– –2H+

t-Bu

t-Bu

SCHEME 26.58

t-Bu

t-Bu

–t-Bu+

H2O t-Bu 96

–2e– –H+

HO t-Bu 97

O 98

Transformation of 2,4,6-tri-tert-butylphenol (96) in presence of water.

nucleophiles, including solvent, electrolyte, additives, or impurities. Scheme 26.57 displays the general procedure of this useful phenol oxidation, and Table 26.7 provides some examples of this transformation. Table 26.7 includes no example for the nucleophilic trapping of the phenoxonium ion by water. Such syntheses of quinones have been carried out by the oxidation of phenols or para-alkoxyphenols in the presence of water [123,127]. The para-alkoxyphenols yield the desired quinone via the generation of the corresponding hemiacetals. Scheme 26.58 shows the transformation of 2,4,6-tri-tertbutylphenol (96) into quinone 98. The reaction was conducted in acetonitrile involving traces of water [128]. The initial electron transfer followed by nucleophilic attack of water gives access to para-quinol 97. Para-alkoxyphenols under comparable conditions lead to the formation of hemiacetals. Extrusion of a tert-butyl cation and subsequent oxidation generates the quinone. A useful application of this transformation is depicted in Scheme 26.59. The conversion of p-substituted phenols with formaldehyde results in an architecture called calix[4]arene, which is widely used in supramolecular chemistry (99) [129]. These structures and their corresponding quinones are interesting as redox-sensitive sensors [130]. The oxidation takes place on platinum anode in acetonitrile involving traces of water with tetrabutylammonium perchlorate as supporting electrolyte. Beneficially, the oxidation is accompanied with the loss of the tert-butyl moieties and due to the cationic charges, a selective conversion of opposite arene moieties is achieved to form calix[4]arenediquinone (100). Electrolysis of para-alkoxyphenols in alcoholic solvents like methanol allows access to synthetically useful quinone monoketals, which can be used as precursors for a rich follow-up chemistry [125,131]. A variety of alkyl substituents in position 2 are tolerated (Scheme 26.60). The electrolytic

© 2016 by Taylor & Francis Group, LLC

1008

Organic Electrochemistry

HO

O

Oxidation Pt, CH3CN O O H HC 3

O

O CH3 H

OH O H3C

O

O CH3 H

O O H 3C

O

O CH3 H

99 OH

O

O O H3C

O

O

O

O O H3C

HO CH3

O CH3

O

100

SCHEME 26.59

Electrochemical preparation of calix[4]arenediquinone (100). OH

O R

R

MeOH, LiClO4 R = H, CH2Me, CH2t-Bu

MeO OMe

OMe

SCHEME 26.60 Synthesis of quinone monoketals.

conversions were carried out in methanol with lithium perchlorate as supporting electrolyte by using platinum as anode material. Good yields up to 85% were reported. Investigations of wastewater degradation at BDD anodes and carbon–PTFE gas diffusion cathodes indicate that oxidation of sulfanilic acid (101) to 1,4-benzoquinone (102) is possible (see Scheme 26.61) [132]. In this case, no direct electron transfer to the sulfanilic acid occurs. Instead, the hydroxyl radicals act as oxidizer. There are two ways to generate hydroxyl radicals for the hydroxylation of sulfanilic acid (see Scheme 26.62). On the one hand, cathodic reduction of oxygen can be used to generate hydrogen peroxide to form the radical by the action of Fe2+ (electro-Fenton). Fe2+ can be regenerated by the reduction of Fe3+ to Fe2+ on the cathodic surface. On the other hand, hydroxyl radicals are obtained by the anodic oxidation of water on BDD. In contrast to the previously described methods, quinones can also be formed by the oxidation of polyfluorinated benzenes involving a 1,4-difluoro situation [133]. Despite the electron-deficient nature of the substrates, the access to such highly fluorinated products is of significant interest, O2 + 2H+ +2e– H2O2 SO3H

Fe2+

+2 OH

–SO42–, –NH4+ NH2 101

SCHEME 26.61

O

OH

OH

–2H2O OH

O 102

Wastewater degradation to form 1,4-benzoquinone (102) from sulfanilic acid (101).

© 2016 by Taylor & Francis Group, LLC

1009

Oxygen-Containing Compounds O2 + 2H+ + 2e–

H2O2

Fe2+ + H2O2

Fe3+ + OH + OH–

H2O

OH + H+ + e–

Fe3+ + e–

Fe2+

SCHEME 26.62 Different ways to produce hydroxyl radicals.

F

F

–2e– 1) TFA/CH2Cl2 NEt3

F

F

2)

F

O F

F

F

F

Yield: 72%

H2O

O 103

F

F

F

F

F

F

F

2)

F

O

F

F

F

F

Yield: 52%

H2O

F

F

SCHEME 26.63

–2e– 1) TFA/CH2Cl2 NEt3

F

O 104

Conversion of 1,4-difluoro arenes into para-quinone.

since the conversion is tremendously easier than the installation of the fluorine moieties in the final product. Scheme 26.63 displays two examples of the electrochemical perfluoro-1,4-benzoquinone synthesis 103, 104. These electrolyses were performed on platinum electrodes in divided cells with methylene chloride and trifluoroacetic acid as solvents and triethylamine as supporting electrolyte. The quinones are afforded upon aqueous work-up. The detailed reaction mechanism of this trifluoroacetoxylation is postulated and plausible for the conditions applied (Scheme 26.64) [133]. Basically, there are two pathways for the generation of cation 107. Path A starts with the oxidation of hexafluorobenzene (105) to cation radical 106 followed by a nucleophilic attack of the trifluoroacetate and another oxidation step to cation 107. Path B Path A F

F

F

F

F

F

–e–

F 105

F

F

F

F

F

OOCCF3 F CF3COO–

F TFA

F

F

F 106

F

OOCCF3 F

F F

F 107

F

F3CCOO

F

–e– Path B

F –

CF3CO2H

–e

–H+

CF3CO2 108

+105

F F

Hydrolysis

F

F

F

F

SCHEME 26.64 Mechanism of quinone formation from 1,4-difluoro arenes.

© 2016 by Taylor & Francis Group, LLC

O

OOCCF3 F

F F O 109

1010

Organic Electrochemistry

NR2

SCHEME 26.65

NHR2H

+ AcOH

+ AcO–

Acetoxylation with solid support bases. OH

O

N t-Bu

t-Bu

O

t-Bu

t-Bu

AcOH/MeCN t-Bu 110

SCHEME 26.66

t-Bu OAc 111

Para-acetoxylation of of 2,4,6-tri-tert-butylphenol (110) with solid support bases.

starts with the oxidation of trifluoroacetic acid to get the trifluoroacetyl radical (108). This radical is powerful enough to attack hexafluorobenzene. The open-shell intermediate is oxidized directly or indirectly, resulting in cation 107. Subsequent nucleophilic attack of trifluoroacetate and hydrolysis accomplishes the final quinone 109. For understanding the mechanistic course of the acetoxylation reaction of hydrocarbons, oxidation potentials of the acetate and the hydrocarbon compounds were determined [107]. In acidic media, the oxidation potential of the acetate ion is higher than the potential of hydrocarbon compounds like naphthalene. Mechanistic investigations of the acetoxylation reaction indicate that this reaction takes place at a potential closer to the potential of hydrocarbon substrate than the one of the acetate. In that case, reaction path A seems to be more realistic. The oxidation potential for the trifluoroacetate is significantly higher than the one for acetate. Consequently, both pathways seem to be viable and most probably occur. Recent investigations of the acetoxylation of phenols deal with the application of solid supported bases [134]. The concept relies on an acid–base reaction between acetic acid and bases supported on silica gel. The immobilized base is able to form acetate required for the reaction course (see Scheme 26.65), but these bases are stable under acidic and electrochemical conditions. Therefore, separation of base after electrolysis is easy and recycling for subsequent runs possible. Beneficially, no further electrolyte has to be added, which generates a significant green aspect onto such conversions. Scheme 26.66 shows a representative example for such a solid supported base acetoxylation reaction. The electrolysis was carried out in an undivided cell equipped with platinum electrodes. Acetoxylation of 2,4,6-tri-tert-butylphenol (110) yields 95% of the para-acetoxylation product 111. In summary, the electrochemical conversion of phenols into the corresponding quinones is a very useful reaction, which can be synthetically exploited on a preparative scale. Different starting materials can be transformed into the desired quinones in good to excellent yields. This particular fact might have an impact on the use of substrates mixture originating from waste streams. In addition, highly reactive quinones can be generated in situ as intermediates for following up sequences. 3. Phenol Coupling Despite the ongoing discussion on which mechanism applies when aryl–aryl bonds are formed, the anodic coupling will definitely start with an initial oxidation step. For phenolic substrates, this will be particularly the case, since the electrolyses are often carried out in basic electrolytes and transformation is dependent on the oxidation potential. Therefore, the term oxidative coupling should be rather used instead of Scholl reaction [135a]. The use of electroorganic methods instead of stoichiometric reagents for the generation of phenol coupling products is of great practical interest, since several of the products are useful building blocks for ligands [135b]. In addition, the moieties formed are also present in naturally occurring products [135b].

© 2016 by Taylor & Francis Group, LLC

1011

Oxygen-Containing Compounds

a. Homocoupling of Phenols As already outlined, phenols exhibit a variety of reaction modes and several reactive positions of the occurring intermediates. The electrolysis conditions applied are crucial and determine the molecular architecture that is obtained. In particular, simple alkyl-substituted phenols show a broad scope of products. When p-cresol (112) is anodically treated at carbon anodes in acetonitrile, the Pummerer ketone scaffold 116 is formed and isolated in 37% yield (Scheme 26.67) [136]. The  Pummerer ketone originates from an ortho–para coupling with a subsequent 1,4-addition. The key for isolating substantial amounts from the complex product mixture is the addition of sodium hydroxide to the electrolyte. Consequently, the substrates for electrolysis are the corresponding phenoxides, which were treated with 1 F (Scheme 26.67). 2,4-Dimethylphenol (114) represents a rather simple substrate but experienced significant attention in the anodic coupling reaction, since the ortho, ortho-coupled product is a 2,2′-biphenol and an important precursor for several catalysts [137]. The treatment on lead dioxide anodes in diluted sulfuric acid provides mostly the hydroxylation product 120 in 44% isolated yield, whereas the 2,2′-biphenol 121 is only obtained in 18% yield (Scheme 26.68). The conversion for the electrolysis is about 94% [127]. Switching with the substrate 114 to basic electrolytes creates a vast variety of anodic products, which underlines the pronounced tendency of 114 for manifold reactions. The anodic transformation can be focused when using, for example, barium hydroxide as supporting electrolyte in methanol (Scheme 26.69) [138]. In accordance with the results obtained with cresol [136], the desired biphenol represents only a minor component. The major is the Pummerer ketone 118 and some follow-up pentacyclic products 122–124. An efficient electrolysis protocol for the synthesis of the diastereomerically pure spiropentacycle 122 was elaborated. The anodic treatment in an undivided cell equipped with platinum electrodes and in the presence of barium hydroxide results in the precipitation of intermediate 125 (Scheme 26.70), which represents a dehydrotetramer of 114 [139]. Since the dehydrotetramer 125 precipitates during the course of electrolysis, the isolation is easy to perform and applicable in OH R3

Carbon anode H3CCN, Et4NBF4

R2

NaOH

R3 R2

O H O

R1

R3

R1

R1

R2 112 R1 = CH3 113 114

116 = 37%

R3 = H

117 = 31%

2= H

R

118 = 20%

R2 = R3 = H

119 = 25%

3 R1 = R3 = CH3

115 R1 = Et

SCHEME 26.67

R2 = R3 = H

R1 = R2 = CH

Formation of Pummerer ketone on carbon anodes in alkaline media. OH

O

HO

PbO2 anode + dil. H2SO4 OH 114

OH

44%

18%

120

121

SCHEME 26.68 Acetoxylation product of 2,4-dimethylphenol (114) on lead dioxide anodes.

© 2016 by Taylor & Francis Group, LLC

1012

Organic Electrochemistry

OH

HO

Pt electrodes 12.5 mA cm–2

O H +

MeOH, base

O

OH

114 32% 118

3% 121

O O O +

+

O

O O

18%

5% 123

122

O O

+

O O

4% 124 (two isomers 1:1)

SCHEME 26.69 Variety of anodic products from 2,4-dimethylphenol using basic conditions. OH

Pt electrodes Ba(OH)2∙8H2O MeOH

O

HO O

O H+ or TiCl4

O O

O

114 52–60% 125

75% 122 TFA, Et3SiH

BF3∙OEt2 –65°C

O O

O

65% 127

83% 128

SCHEME 26.70 Dehydrotetramer 125 of 2,4-dimethylphenol (114) as starting material for different polycyclic architectures

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Oxygen-Containing Compounds

1013

multigram scale. Removal of one equivalent 2,4-dimethylphenol induces a stereospecific rearrangement to the spiropentacycle 122. The latter step can be effected thermally or by acidic treatment. 125 represents a key intermediate for the selective synthesis of a diversity of polycyclic architectures, whose selective formation from 114 can be controlled by the reaction conditions [140]. More than 14 scaffolds, including, for example, propellanes and dibenzofuranes, are more or less exclusively obtained from this electrochemically generated precursor [141]. Further formations of unique architectures by anodic treatment are treated in Chapter 18. The direct and selective conversion of 2,4-dimethylphenol to the desired 2,2′-biphenol 121 can be synthetically achieved when using specific carbon electrodes and suitable electrolytes. When employing BDD electrodes, a selectivity for the desired 121 compared to the Pummerer ketone derivative is better than 18:1 [142]. The electrolysis is conducted in almost neat phenol at 70°C and about 11% water in the electrolyte system. Although a supporting electrolyte is present, water is necessary to provide sufficient conductivity. Using solvent-based electrolytes results mostly in an electrochemical incineration [143]. The electrolysis only provides 2,2′-biphenol 121 in 56% yield and with 29% current efficiency. Overoxidation of the desired product limits the yield in a batch-type electrolysis. Performing only partial conversion of phenol 114 with approximately 0.4 F avoids significant oligomer formation, and the nonconverted substrate can be redistilled by short-path distillation [144]. The unique properties of BDD can be imitated by glassy carbon and trifluoroacetic acid as electrolyte system. The electrolysis of 114 to 2,2′-biphenol 121 is accomplished in 64% yield on recovered starting material. The current efficiency is 53%, and the corresponding Pummerer ketone is not anymore detected [145]. The addition of 1,1,1,3,3,3-hexafluoroisopropanol to the electrolyte allows a control of the reactivity when using BDD anodes [12]. The electrolysis of 114 in an undivided cell provides 47% isolated yield of 2,2′-biphenol 121. This protocol seems to be general for the electrolytic generation of symmetric biphenols in acceptable yields [146]. The successful tuning in reactivity is underlined in the anodic treatment of sesamol to the corresponding 5,5′-linked dehydrodimer 130. The isolated yield is 74%, whereas other electrochemical conditions only provide tar-type residues [12]. For less activated phenols, it is not necessarily the more productive protocol. 2-Bromo-p-cresol or 2-fluoro-p-cresol is more efficiently coupled at a graphite anode with trifluoroacetic acid as electrolyte component. Although the conversion of 2-fluoro-p-cresol is rather moderate, it represents the first direct anodic coupling of a fluorophenol. Methyl-triethylammonium methylsulfate (MTES) served as inexpensive electrolyte and gave the best results among all tested supporting electrolytes [145] (Table 26.8). The coupling of guaiacol derivatives 137, 139, and 141 are best performed at BDD anodes with 1,1,1,3,3,3-hexafluoroisopropanol, since this combination renders the best isolated yields [147]. Remarkably, the anodic treatment of 4-methylguaiacol yields selectively the ortho–meta coupling product 136. This indicates that an intermediate phenoxyl radical attacks a closed shell reaction partner on its most electron-rich position. Consequently, other substitution patterns, for example, 2,4-dimethoxyphenol, yield the expected 2,2′-biphenolic product 138 in 45% isolated yield [147]. Steric as well as electronic reasons direct the course of electrolysis and provide the biphenols 140 and 142, in 44% and 83% isolated yield, respectively. The undivided electrolysis is easy to perform and generates these products in exclusive selectivity [147]. The electroorganic coupling of Schiff bases derived from ortho-vanillin results in the dehydrodimers 144 in good to excellent yields [148]. The use of glassy carbon as anode and acetonitrile with perchlorate as electrolyte turned out to be suitable. In general, 2,6-dialkyl-substituted phenols are prone to the formation of quinoide dehydrodimers. However, electrolysis at low voltage in a divided cell with frequent reversal of polarity provides the 4,4′-biphenol in excellent yield [149]. The anodic conversion of eugenol in basic electrolyte at platinum anodes to the 2,2′-biphenol 148 is reported to be a quantitative process [150]. When less electron-rich phenols, for example, 2-acetyl-p-cresol, are treated under analogue conditions, only a moderate amount of the 2,2′-biphenol 150 is obtained [151]. Having an additional methoxy group in the substrate for electrolysis, for example, in vanillin, is rewarded by significant better results of 65% yield [152].

© 2016 by Taylor & Francis Group, LLC

1014

Organic Electrochemistry

TAbLE 26.8 Homocoupling of Different Phenols Substrate

Electrolysis Conditions

Product

BDD anode, MTES, 11% H2O BDD anode, MTES, HFIP Graphite, MTES, TFA

OH

yield (%)

References

56

[142]

47 64

[12] [145]

74

[12]

30 76

[12] [145]

19 21

[12] [145]

33 57

[147] [145]

45

[147]

44

[147]

a

OH

OH 114

121 OH

O

BDD anode, MTES, HFIP

O O OH

O 129 OH

O O

130 Br OH

BDD anode, MTES, HFIP Graphite, MTES, TFA

OH Br Br OH

131

132 F OH

BDD anode, MTES, HFIP Graphite, MTES, TFA

OH F

F 133

OH 134

O OH

BDD anode, MTES, HFIP Graphite, MTES, TFA

O OH OH

135

O 136 BDD, MTES, HFIP

O

O

OH

OH O

O

O 137

OH 138 BDD, MTES, HFIP

OH O

O

O

OH O

O O

139

OH

140 O

(Continued )

© 2016 by Taylor & Francis Group, LLC

1015

Oxygen-Containing Compounds

TAbLE 26.8 (Continued) Homocoupling of Different Phenols Substrate

Electrolysis Conditions BDD, MTES, HFIP

OH

Product HO

yield (%)

O

O

O

[147]

41–100

[148]

93

[149]

100

[150]

26

[151]

O

O O

141

References

83

OH

142 Glassy carbon, NaClO4, H3CCN

OH O

O HO

NR RN

O

R = alkyl, benzyl 143

OH 144 NR OH

tBu

tBu

Pt, LiClO4, MeOH/CH2Cl2, divided cell

tBu HO

tBu tBu 145 OH 146

tBu OH

Pt, LiClO4, NaOH, MeOH

O OH O

O OH 148 147 O

OH

Pt, NaOH, MeOH/H2O OH

O 149

O

OH 150 (Continued )

© 2016 by Taylor & Francis Group, LLC

1016

Organic Electrochemistry

TAbLE 26.8 (Continued) Homocoupling of Different Phenols Substrate

Electrolysis Conditions

Product

Pt, NaClO4, Et4NOH, H3CCN

OH O

yield (%) 65

O

References [152]

OH O

O OH O 151 152

O

MTES, Et3NCH3+ −O3SOCH3; HFIP, 1,1,1,3,3,3-hexafluoroisopropanol; TFA, trifluoroacetic acid. a Based on recovered starting material.

In order to circumvent the nondesired pathways upon the oxidation of phenols and to gain selectivity, the introduction of a template at the phenolic oxygen atoms is a common and actual topic [153]. For employing conventional reagents for the oxidative transformation, covalently tethered phenolates were successfully applied. Several elements, such as titanium, zirconium, silicon, and phosphorus, have been studied as oxophilic templates for enolates or phenolates in oxidative coupling processes [154]. Silicon-tethered substrates turned out to exhibit the best synthetic practicability [155]. For the anodic treatment, boron is the element of choice, since the limited coordination number, stability, and hydrolytic cleavage are ensured. Moreover, 153 can be prepared from the corresponding phenols in a two-step protocol or a one-pot procedure [156]. As sodium base, either a piece of sodium metal is added after removal of water or sodium methanolate is used. Moreover, these tetraphenoxy borates 153 are anionic and serve as substrates as well as supporting electrolytes. Consequently, phenoxy borates are ideal substrates in terms of tethered phenolates [157]. The oxidation potential corresponds to the one of the individual phenolates. The anodic treatment in acetonitrile results in 154, which can be hydrolyzed with hot water. The broad scope of this transformation is demonstrated in more than 20 examples [157]. This electroorganic approach can be applied to a variety of phenolic substrates and is scalable to a multi-kg range [158]. The electrolysis can be performed on platinum or graphite anodes. The overall yields are acceptable to good with a superb product quality. The scope is limited by the necessity of a substituent in position 4 of the phenolic substrate and the ligand exchange during electrolysis. Because of the latter the electrolysis is best carried out at elevated temperature and a chiral dummy ligand cannot be exploited for stereo induction (Scheme 26.71). b. Cross Coupling with Phenols The direct electrochemical cross coupling faces the challenge that usually the individual oxidation potential of the components is a strong selectivity directing argument. This will result in a strong preference for homo-coupling products 158 and 159, whereas the desired cross coupling product 160 is often only obtained in traces. Even when the more stable component is used in excess, the mixed biaryl is obtained in far less amounts than statistics would predict (Scheme 26.72). Yoshida et al. circumvented this dilemma with the cation-pool method and established an elegant electroorganic access to unsymmetrical biaryls [159]. So far, this method cannot be extended to biphenols and is not scalable, since electrolysis at very low temperature and expensive supporting electrolytes is required. However, the spatially and timely separated oxidation of one phenol to a quinone ketal by hypervalent iodine reagents and subsequent conversion with another phenol derivative provides mixed biphenols [160]. The strong electrophilic conditions, the two-step sequence,

© 2016 by Taylor & Francis Group, LLC

1017

Oxygen-Containing Compounds OH

R B(OH)3, –H2O

4

R

Then Na or Na base

Na B

Anode H3CCN

O 4

153 Na+

R O O

– B

R Hydrolysis

O O

OH

R

2

R

OH R

R 154

OH

OH

OH

SCHEME 26.71

OH

HO

85%

66%

48%

155

156

157

Selective biphenol preparation via boron templates. H

H

Oxidation R

OH

+

R

R΄ +

R΄ R΄

R 159

158 R +

R΄ 160

SCHEME 26.72 Anodic phenol-phenol cross-coupling.

and excess of oxidizer limit the scope of this conventional phenol cross coupling. A direct arylation of phenols was achieved by using guaiacol derivatives and electron-rich arene components [13]. A rationale is outlined in Scheme 26.73. At a BDD anode, the phenolic substrate is oxidized, and after extrusion of a proton, a phenoxyl radical 161 is generated. The electrophilic behavior of this intermediate allows conversion with electron-rich arenes. The second oxidation step of 164 can occur directly or indirectly and accomplishes the mixed biaryl. Because of the two oxidation steps, the process is highly sensitive to the current density applied. Typically low current density in the range of 2–6 mA cm−2 is employed. It turned out to be an electrolyte-dominated process, and the presence of methanol or water in 1,1,1,3,3,3-hexafluoroisopropanol (HFIP) is crucial for a good yield and unique selectivity for the cross coupling product [161,162]. More examples about the coupling of phenols with nucleophilic partners are provided in Chapter 18, whereas this section is dealing with the phenol–phenol cross coupling. Since the found selectivity for the cross coupling can only be explained by the formation of strong solvates, this transformation was extended to the phenol–phenol cross coupling. Dependent

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1018

Organic Electrochemistry + OH R A

H+

O

O

N R΄

R

R O

162

161

D E

O

OH

H

R΄

R΄ R

R

H+

H

H

164

163

OH

R΄

R

SCHEME 26.73 Mechanism for the anodic phenol-phenol cross-coupling.

on the substitution pattern of the phenol, the nucleophilicity and oxidation potential can be decoupled in a certain range. The more electron-rich component experiences tight solvation and the nucleophilicity is strongly decreased, whereas the less oxidizable phenol can act as nucleophile and starts the cross coupling sequence. Besides the highly fluorinated alcohol HFIP, other solvents like formic acid are also suitable for the electroorganic cross coupling [163]. Both solvents have in common a very strong hydrogen bonding capability. The selectivity for the phenol–phenol cross coupling product is in both solvent mixtures higher than 100:1 (see Table 26.9). Because the anodic stability of formic acid is not given, the isolated yields are significantly lower compared to the electrolyses performed in HFIP/MeOH mixtures [164]. The scope for the formation of the unsymmetrical 2,2′-biphenols covers products that combine electronically and sterically different moieties (Scheme 26.74). The example of the solvent-directed cross coupling underlines that this particular field is emerging and similar advances can be expected for other substrates in near future. The full potential of electrosynthesis will be demonstrated when several steps of conventional synthesis can be short-cut, because different pathways could be employed.

B. REDUCTION 1. Electroreductive Hydroxylation of Aromatic Compounds One-step synthesis of phenol from benzene can be used as an attractive alternative for the Hock process, which generates phenol from cumene on an industrial scale. The electrochemical synthesis of phenol uses the reductive regeneration of oxidizing agents like Fenton’s reagent or generates H2O2 reductively from molecular oxygen for the chemical oxidation of benzene. Advantages of this approach are the conversion to phenol at ambient temperature and pressure compared to the Hock process, wherein high temperatures and high partial oxygen pressure are required. In addition, stoichiometric side products, like acetone, are avoided. When using Fenton’s reagent, the electrochemical reduction usually regenerates Fe2+ from Fe3+, which is consumed due to side reactions.

© 2016 by Taylor & Francis Group, LLC

1019

Oxygen-Containing Compounds

TAbLE 26.9 Anodic Cross Coupling of Phenols with Different Solvent Mixtures yield (%) Product

HFIP/MeOH

HCOOH/MeOH

References

50

28

[164]

36

13

[164]

61

45

[164]

63

34

[164]

OH O OH

OH O OH

OH O OH

O OH

OH

MTES, Et3NCH3+ −O3SOCH3; HFIP, 1,1,1,3,3,3-hexafluoroisopropanol; TFA, trifluoroacetic acid.

OH

OH H

R

H +

R΄

BDD anode MTES Solvent mixture

OH

R OH

R΄

Selectivity: >100:1

SCHEME 26.74

Direct anodic phenol-phenol cross-coupling reaction.

© 2016 by Taylor & Francis Group, LLC

1020

Organic Electrochemistry Cathode

H2O2

+

Fe3+

Fe2+

Fe2+

Fe3+

+ HO–

H

H +

R

+

OH

OH

R

OH

OH H

OH +

R

With

SCHEME 26.75

R=H

Fe3+

Fe2+

+

H+

+

R

64%

R = Cl

18% (o:p = 60:40)

R=F

80% (o:p = 85:15)

R = CN

21% (o:m:p = 27:18:55)

Usage of H2O2

Cathodic hydroxylation of benzene derivatives using Fenton’s reagent.

The Fe2+ ions react with hydrogen peroxide to provide hydroxyl radicals. These hydroxyl radicals can oxidize benzene and analogue substrates to phenols (see Scheme 26.75) [165]. The displayed yields in Scheme 26.75 were obtained by employing H2O2, which was not electrochemically generated. When adding copper salts to Fenton’s reagent, oxidation takes place more easily, since Cu2+ is a better oxidizer for carbon-centered radicals than Fe3+ [165,166]. Regeneration of resulting Cu+ is achieved by redox reaction with Fe3+. With this method, various benzene species could be oxidized to their corresponding phenols in moderate-to-good yields (see Scheme 26.75). To prevent further functionalization to hydroquinone or catechol, which represents a general problem in this electroreductive synthesis of phenols, continuous extraction of the phenolic products is a viable strategy [165]. In this conversion, hydroxyl radicals, generated from hydrogen peroxide, act as oxidizer. Direct synthesis of phenol from benzene and oxygen has been achieved in trifluoromethanesulfonic acid as electrolyte on a graphite cathode with current efficiencies of 22–41% [167]. The postulated mechanism involves an electrochemical reduction of oxygen to H3O2+, which is capable of oxidizing benzene (Scheme 26.76). The combination of both methods, reductive generation of hydrogen peroxide in the presence and absence of Fenton’s reagent, was carried out in 0.1 M H2SO4 on mercury, lead, copper, and silver electrodes. In the presence of Fenton’s reagent, hydroxylation of benzene to phenol, hydroquinone, and catechol and the hydroxylation of phenol to hydroquinone and catechol are considerably accelerated compared to the same reaction without Fenton’s reagent [168]. A similar approach was investigated by the oxidation of benzene in 0.1 N H2SO4/acetonitrile (9:1) with Cu+ and oxygen. Consumed Cu+ was regenerated from Cu2+ electrochemically [166]. Hibino et al. described the cathodic formation of phenol from benzene at vanadium oxide–based electrodes under direct and alternating current conditions with high current efficiencies and selectivities. Although the only noticeable side product mentioned is CO2, the selectivity of the reaction is questionable, because the electrochemical reaction was carried out only to low conversions and possible side products may occur on the synthetic scale [169,170]. Another approach for electroreductive synthesis of substituted phenols was investigated by the reduction of nitrobenzene compounds in voltammetric studies and preparative scale electrolysis.

© 2016 by Taylor & Francis Group, LLC

1021

Oxygen-Containing Compounds +

OH2 – – – – – – – – – – – – – – – – – – – –

OH2 OX H3O2+ or "OH+"

OX

RED CF3SO3H

O2

Cathode

SCHEME 26.76

+

None

Direct synthesis of phenol from benzene and oxygen. NO2

NH2

NHOH

Cu(Hg) cathode 4e–, 4H+

H+

20% H2SO4(aq) Cu(Hg) cathode 2e–, 2H+

20% H2SO4(aq)

OH 76.9%

NH2

18.8%

SCHEME 26.77

Reduction of nitrobenzene to p-aminophenol and aniline.

Reduction of nitrobenzene with an amalgamated copper cathode in 20% aqueous H2SO4 yields almost 77% p-aminophenol after Bamberger rearrangement and 18.8% aniline as a side product [171] (Scheme 26.77). 2. Electroreductive birch-Type Reaction Electrochemical Birch-type reactions (see Scheme 26.78) of different aromatic systems have become a synthetically important alternative to the classic Birch reduction using alkali metal in liquid ammonia, even for presumably more electron rich and consequently more difficult to reduce systems like methoxyaromatics and phenol [172]. It is noteworthy that only after the protection of phenols as ether a practical method is established. Generally, the cathodic Birch reduction has been carried out successfully in various solvents, such as liquid ammonia [173], methylamine [174], ethylenediamine [175], hexamethylphosphoramide [176], aqueous diglyme [177], anhydrous THF [178], and water [172]. The advantages of a cathodic Birch reduction compared to the classic protocol are mostly due to the possibility to avoid working with large quantities of alkali metal and a much simpler experimental procedure. For instance, it is possible to avoid liquid ammonia, which is a strongly

© 2016 by Taylor & Francis Group, LLC

1022

Organic Electrochemistry O

O Cathode +2H+

R

SCHEME 26.78

R

Electrochemical Birch-type reaction.

(C4H9)4N+

Hg cathode

(C4H9)4N (Hg)n

O (C4H9)N (Hg)n +

O

R

O +

H2O

O +

(C4H9)4N (Hg)n

SCHEME 26.79

+

R

O

(C4H9)4N+ + nHg

O +

R

– OH

+

R

O R

+ (C4H9)4N+ + nHg

R

–

R

–

O

H2O

R

+

– OH

Electrochemical Birch-reaction at a mercury cathode.

basic solvent and has to be constantly kept below −33°C; also solubility of the reactants becomes a minor issue due to the larger variety of solvents, which can be employed. In particular, for upscaling, the electroorganic approach provides several beneficial safety aspects. The mechanism of a cathodic Birch reduction strongly depends on the used electrolyte and electrode system. Studies on the electroreduction of methoxybenzene, 1,2,3,4-tetrahydro-6-methoxynaphthalene, and steroidal systems using aqueous tetrabutylammonium hydroxide and a mercury cathode report a suggested single-electron transfer from an electrogenerated tetrabutylammonium amalgam [172] (Scheme 26.79). In contrast to this postulated mechanism, electrolyte systems based on lithium salts in combination with magnesium electrodes in THF suggest an electron transfer from electrogenerated lithium species similar to the conventional Birch reduction [178]. In all cases, the product yield is highly sensitive to solvents, electrode materials, and temperature applied. Besides various studies on substituted benzene and naphthalene systems, further investigations were carried out on cathodic Birch reduction of methoxy-substituted steroids (see Table 26.10). Cathodic Birch reduction on steroids became an important synthetic alternative to the classic procedure of the Birch reduction, since many substrates exhibit a challenging solubility in liquid ammonia [180]. Kashimura et al. described a method to incorporate deuterium in position 1 and 4 of aromatic systems when electrochemical Birch reduction is performed in anhydrous THF in the presence of 10 equiv. t-BuOD instead of t-BuOH as proton source [178]. The latter was not performed directly on phenols as substrates.

© 2016 by Taylor & Francis Group, LLC

1023

Oxygen-Containing Compounds

TAbLE 26.10 Electrochemical birch Reaction for Different Substrates Electrolyte (Cathode)

Substrate

H2O–Bu4NOH (Hg)

O

yield (%)

Product O

H2O–Bu4NOH (Hg)

References

80

[172]

85

[172]

91

[178]

80

[178]

93

[179]

92

[179]

70

[173]

O

O THF–LiClO4 (Mg)

O

O THF–LiClO4 (Mg)

O O

O

O O

O

MeOH/MeCN (20/3)–Et4NCl (Pb/Pt)

O

O O

O

+

O

O O O

THF/MeOH– Et4NCl (Pb/Pt)

O

O

O +

O

O OH

Liq. NH3–NaCl/ NaI (Mg/Al)

OH

(Continued )

© 2016 by Taylor & Francis Group, LLC

1024

Organic Electrochemistry

TAbLE 26.10 (Continued) Electrochemical birch Reaction for Different Substrates Electrolyte (Cathode)

Substrate

THF/ H2O–Bu4NBF4

OH

O

yield (%)

References

OH

n.d.a

[180]

OH

n.d.a

Product

O +

O OH R

H2O–Bu4NOH (Hg)

O

[172]

R

O

R = H or C2H5 a

90

OH

R = H or C2H5

Ratio of depicted products 10:90, but no clear data of isolated yield are given.

3. Electroreductive Hydrogenation and Hydrodeoxygenation of Phenolic Moieties Electrocatalytic hydrogenation of unsaturated molecules is a process wherein acidic solvents are used as a proton source to generate chemisorbed hydrogen at the cathode. Further hydrogenation reaction mostly proceeds similar to catalytic hydrogenation reactions. This method avoids the need for an external, fossil-based hydrogen supply and therefore allows a mild reaction at ambient conditions. This implies less safety aspects compared to typical hydrogenation reactions. Electrocatalytic hydrogenation of organic molecules is of current and high interest in the context of green chemistry [181]. Electroreductive conversion of phenol to cyclohexanol has been reported in a 2 N aqueous HClO4 – Et4NBr system at a platinum cathode in 82.7% yield with a current efficiency of 69.5%. Under similar conditions, 2-naphthol could be converted to trans-decalol in 70% yield (see Scheme 26.80) [182]. OH

OH Pt cathode 2 N HClO4 in H2O Et4NBr 82% Pt cathode

HO

2 N HClO4 in H2O Et4NBr

HO 70%

SCHEME 26.80 Electrochemical hydrogenation of phenol and naphthol.

© 2016 by Taylor & Francis Group, LLC

1025

Oxygen-Containing Compounds OH

OH Pt cathode H2O/H2SO4 100% c.e.

X

X

X = H, OH, OMe, Me, COOH, NH3+ NH2CH3+ , CN, Et, tert-Bu

SCHEME 26.81 Electroreductive hydrogenation of phenol derivatives.

To reduce phenol to cyclohexanol with a reasonable current efficiency, the reaction has to be carried out at low current densities (0.67 mA cm−2) [166]. Therefore, large amounts of solid platinum foil have to be employed to ensure a high reaction rate. In order to avoid the use of high amounts of platinum, the electrocatalytic hydrogenation of phenol to cyclohexanol on highly dispersed platinum electrodes has been developed [183]. Current efficiencies of 85% could be achieved on electrodes with a platinum loading of 2%. Due to the small amount of platinum, which is used in highly dispersed platinum electrodes, the cost in the fabrication of these electrodes is much less than bare or platinized platinum [183]. Recent investigations are carried out on finding an effective support material/catalyst pair for catalytic nanoaggregates, since the support material of the electrode has been found to play an important role in the effectiveness of electrocatalytic hydrogenation reactions [181]. Electrocatalytic hydrogenation can be used for stabilizing biooil by hydrogenation and hydrodeoxygenation of phenolic compounds. Therefore, the reduction of three model substrates, phenol, 2-methoxyphenol, and 2,6-dimethoxyphenol, has been investigated. The electrocatalytic species was generated by ruthenium on activated carbon cloth. The conversion was achieved at mild temperatures and ambient pressure [184]. Another mechanistic procedure has been found in the reduction of phenol and anisole nuclei when reduction to cyclohexanol and methoxycyclohexane was carried out in an EtOH–HMPA–LiCl system at an aluminum cathode in 45.4–51.7% yield. In this protocol, solvated electrons as reducing agent were proposed [185]. Therefore, the reaction mechanism is more similar to cathodic Birch reductions. Sasaki et al. describe the electroreductive hydrogenation of the aromatic core of different phenols [166]. Unfortunately, this represents only a cyclic voltammetric study without the isolation of any product. The synthetic utility is consequently questionable (Scheme 26.81). 4. Deoxygenation of Phenolic Compounds Deoxygenation reactions are still one of the major challenges in organic electrochemistry. Indirect deoxygenation of phenolic compounds by electrochemical reductive processes can be achieved when the hydroxyl group is converted into a good leaving group like phosphate esters [186] or tetrazolium ethers [187]. The electrochemical reduction of aryl diethyl phosphates has been carried out on a lead cathode in DMF/p-toluenesulfonate on various substrates giving isolated products in 43–73% yield. The initiation step of this deoxygenation reaction is found to be a one-electron transfer to the aromatic nucleus [186]. Although this procedure is not applicable for all kinds of phenolic compounds, for example, conjugated double bonds in side chains and half of the naphthalene nucleus of 2-naphthyl diethyl phosphate were also hydrogenated along this pathway, it gives good yields especially for methoxy-substituted electron-rich phenols (see Scheme 26.82). A conventional method for the removal of phenolic hydroxyl groups is the conversion into tetrazol-5-yl ethers followed by catalytic hydrogenation in the presence of a noble metal catalyst [187]. The electrochemical pathway has been investigated at a Hg pool cathode in DMF–H2O (95:5 v/v%). The proposed mechanism describes a required one-electron transfer to the phenolic core for a successful cleavage of the aryl-oxygen bond. As anticipated, phenols exhibiting electron-withdrawing groups like p-cyanophenol are more selectively converted. If the substrate is based on electron-rich phenol derivatives, the first electron transfer is addressed to the tetrazole moiety, and therefore, bond cleavage happens in the way that phenolates can be recovered as a side reaction or even as the only conversion product (see Scheme 26.83) [187].

© 2016 by Taylor & Francis Group, LLC

1026

Organic Electrochemistry O

O P

O

O O

O

Pb cathode divided cell

O

O

DMF Et4NOTs 67%

SCHEME 26.82 Cathodic deoxygenation of aryl diethyl phosphates. N=N N N Ph

H

OH

Hg cathode

O

+ Et4NBr DMF–H2O (5 vol%)

R

R

R

R = CN

53% (isolated yield)

0%

R = OMe

30% g.l.c analysis

30% g.l.c analysis

SCHEME 26.83 Deoxygenation of phenols via derivatization with tetrazoles. OAc

OAc Hg cathode Et4NI DMF

O

OH

R

SCHEME 26.84

75%

R

Electrochemical cleavage of a carbon oxygen bond on phenol ethers at a mercury cathode.

Cleavage of a carbon–oxygen bond on phenol ethers has been carried out in DMF–Et4NI on a mercury cathode in chemical yields of 75% for a vitamin K precursor (see Scheme 26.84) [188]. 5. Reduction of Quinones Quinone-based redox couples have been investigated intensively because of their chemical and biological importance, for example, as prototypical reversible redox systems in chemical reactions, in quinoenzymes, and anticancer drugs. The most important function of quinones is the reversible electron and proton transfer. The related thermodynamic parameters of interest are the standard redox potential E° and the pK values. The standard redox potential E° of different quinone systems strongly depends on the quinone series, for example, benzoquinone, naphthoquinone, or anthraquinone, the substituents of the quinone ring, and the stabilization of the reduced species. Solvent polarity, pH, intra- and intermolecular hydrogen bonding, and ion-pair formation play an important role in the stabilization of the reduced species [188]. In aprotic solvents, quinones are reduced in two successive one-electron steps to an anion radical Q• − and the hydroquinone dianion Q2− (see Scheme 26.85) [189]. Q + e– Q

SCHEME 26.85

–

+ e–

Q

–

Q2–

Reduction sequence of quinones in aprotic solvents.

© 2016 by Taylor & Francis Group, LLC

1027

Oxygen-Containing Compounds O

OH + 2e– + 2H+

O

OH

(a) O–

O + 2e–

O–

O (b)

SCHEME 26.86 Possible reduction products of quinones in acidic (a) and alkaline media (b). O

O –e–, –H+ Electrooxidation

OH

O

Cytotoxic ROS O22–

O

O

O

O

O

+2e–, +2H+

+2e–, +2H+

n

OH O2 OH

O n

Cathode

SCHEME 26.87 Electrochemical reduction of oxygen via quinone polymer films on cathodes.

In acidic solutions, the reduction is a single-step two-electron two-proton process (see Scheme 26.86a), while in alkaline pH, the reduction does not involve protons and is only a two-electron reduction (see Scheme 26.86b) [190]. At neutral pH, reduction is either a one-proton two-electron or only two-electron process without the participation of protons [190]. The electrochemical reduction of quinones can be used synthetically to reduce oxygen as the resulting hydroquinones work as reducers for oxygen. It was possible to use quinones as a polymer film on electrodes capable of serving as a controllable source of reactive oxygen species (see Scheme 26.87) [191]. Electroreductive coupling of 3,4-estrone-o-quinone with adenine has been achieved in a DMF– LiClO4 at a platinum cathode with a chemical yield of 14.1% (see Scheme 26.88) [192]. Polyether-bridged quinones (see Scheme 26.89) can be used as crown ethers equipped with redox systems and are known for complexation with lithium, sodium, and potassium cations [193]. Coupling between the complexation behavior and redox properties of these systems has been proved by cyclic voltammetry [193].

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1028

Organic Electrochemistry N O

N

H2N

NH2

O N

N

Pt cathode

+ N H

N

N

NH

DMF LiClO4

O O

HO OH

SCHEME 26.88

Electroreductive coupling of 3,4-estrone-o-quinone with adenine.

O O

n

O O

O O

O

n

O OH

O

[H] [O] R

R

R

R OH

O R = H, Me n = 1–4

SCHEME 26.89 Polyether-bridged quinones as redox sensitive crown ethers. O

OH Hg cathode EtOH/H2SO4

O

OH

+ OH2

OH Hg cathode EtOH/H2SO4

OH

SCHEME 26.90

Electrochemical reduction of anthrahydroquinone to anthrol and anthrone.

The reversible reduction of anthraquinone to anthrahydroquinone has been carried out in EtOH (50%)/H2SO4 (50%) system on a mercury electrode with further reduction of protonated anthrahydroquinone to anthrol and isomerization to anthrone (see Scheme 26.90) [194]. An intramolecular cyclization product was isolated in 14% yield after the reduction of a bridged benzoquinone in anhydrous CH2Cl2–Bu4NBF4 at a mercury cathode and stabilization of the reaction product by acylation (see Scheme 26.91) [195].

O

1) Hg cathode CH2Cl2

O O

AcO AcO

Bu4NBF4

O

2) Ac2O 14%

O

SCHEME 26.91

OAc

Intramolecular cyclization of tethered benzoquinones.

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Oxygen-Containing Compounds

1029

In summary, the cathodic treatments of phenols or to obtain phenols are viable electrosynthetic pathways that are still far underrepresented. The reason for that might be the limited number of suitable cathode materials, which cause also severe toxicity concerns. Consequently, future research in that particular field is required.

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45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. 61. 62. 63. 64. 65. 66. 67. 68. 69. 70. 71. 72. 73. 74. 75. 76. 77. 78. 79. 80. 81. 82. 83. 84. 85. 86. 87. 88. 89. 90. 91. 92.

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Oxygen-Containing Compounds 93. 94. 95. 96. 97. 98. 99. 100. 101. 102. 103. 104. 105. 106. 107. 108. 109. 110. 111. 112. 113. 114. 115. 116. 117. 118. 119. 120. 121. 122. 123. 124. 125. 126. 127. 128. 129. 130. 131. 132. 133. 134. 135.

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Organic Electrochemistry

137. (a) Alexakis, A.; Polet, D.; Rosset, S.; March, S. J. Org. Chem. 2004, 69, 5660–5667; (b) Li, K.; Alexakis, A. Angew. Chem., Int. Ed. 2006, 45, 7600–7603; (c) Falciola, C. A.; Alexakis, A. Angew. Chem., Int. Ed. 2007, 46, 2619–2622. 138. Malkowsky, I. M.; Rommel, C. E.; Wedeking, K.; Fröhlich, R.; Bergander, K.; Nieger, M.; Quaiser, C.; Griesbach, U.; Pütter, H.; Waldvogel, S. R. Eur. J. Org. Chem. 2006, 241–245. 139. Barjau, J.; Königs, P.; Kataeva, O.; Waldvogel, S. R. Synlett 2008, 2309–2311. 140. Barjau, J.; Schnakenburg, G.; Waldvogel, S. R. Angew. Chem., Int. Ed. 2011, 50, 1415–1419. 141. (a) Barjau, J.; Schnakenburg, G.; Waldvogel, S. R. Synthesis 2011, 2054–2061; (b) Barjau, J.; Fleischhauer, J.; Schnakenburg, G.; Waldvogel, S. R. Chem. Eur. J. 2011, 17, 14785–14791. 142. Malkowsky, I. M.; Griesbach, U.; Pütter, H.; Waldvogel, S. R. Eur. J. Org Chem. 2006, 4569–4572. 143. Griesbach, U.; Malkowsky, I. M.; Waldvogel, S. R. Green electroorganic synthesis using BDD electrodes. In Electrochemistry for the Environment. Comninellis, C.; Chen, G., eds. Springer, Berlin, Germany, 2009, pp. 125–141. 144. Malkowsky, I. M.; Waldvogel, S. R.; Pütter, H.; Griesbach, U. PCT Int. Appl. 2006, WO 2006077204 A2 20060727. 145. Kirste, A.; Hayashi, S.; Schnakenburg, G.; Malkowsky, I. M.; Stecker, F.; Fischer, A.; Fuchigami, T.; Waldvogel, S. R. Chem. Eur. J. 2011, 17, 14164–14169. 146. Fischer, A.; Malkowsky, I. M.; Stecker, F.; Waldvogel, S.; Kirste A. PCT Int. Appl. 2010, WO 2010023258 A1 20100304. 147. Kirste, A. Schnakenburg, G. Waldvogel, S. R. Org. Lett. 2011, 13, 3126–3129. 148. Ohmori, H.; Matsumoto, A.; Masui, M.; Sayo, H. J. Electrochem. Soc. 1977, 124, 1849–1854. 149. Torii, S.; Dhimane, A.-L.; Araki, Y.; Inokuchi, T. Tetrahedron Lett. 1989, 30, 2105–2108. 150. Iguchi, M.; Nishiyama, A.; Terada, Y.; Yamamura, S. Chem. Lett. 1978, 4, 451–454. 151. Johnston, K. M. Tetrahedron Lett. 1967, 8, 841–844. 152. Vermillion Jr., F. J.; Pearl, I. A. J. Electrochem. Soc. 1964, 111, 1392–1400. 153. Masters, K.-S.; Bräse, S. Angew. Chem., Int. Ed. 2013, 52, 866–869. 154. (a) Schmittel, M.; Burghart, A; Malisch, W.; Reising, J.; Söllner, R. J. Org. Chem. 1998, 63, 396–400; (b) Schmittel, M.; Haeuseler, A. J. Organomet. Chem. 2002, 661, 169–179. 155. (a) Takada, T.; Arisawa, M.; Gyoten, M.; Hamada, R.; Tohma, H.; Kita, Y. J. Org. Chem. 1998, 63, 7698–7706; (b) Schmittel, M.; Haeuseler, A. Z. Naturforsch. B 2003, 58, 211–216. 156. Malkowsky, I. M.; Fröhlich, R.; Griesbach, U.; Pütter, H.; Waldvogel, S. R. Eur. J. Inorg. Chem. 2006, 1690–1697. 157. Malkowsky, I. M.; Rommel, C. E.; Fröhlich, R.; Griesbach, U.; Pütter, H.; Waldvogel, S. R. Chem. Eur. J. 2006, 12, 7482–7488. 158. Rommel, C. E.; Malkowsky, I.; Waldvogel, S.; Pütter, H.; Griesbach U. PCT Int. Appl. 2005, WO 2005075709 A2 20050818U. 159. Morofuji, T.; Shimizu, A.; Yoshida, J.-i. Angew. Chem., Int. Ed. 2012, 51, 7259–7262. 160. Parumala, S. K. R.; Peddinti, R. K. Org. Lett. 2013, 15, 3546–3549. 161. Kirste, A.; Elsler, B.; Schnakenburg, G.; Waldvogel, S. R. J. Am. Chem. Soc. 2012, 134, 3571–3576. 162. Elsler, B.; Schollmeyer, D.; Waldvogel, S. R. Faraday Discuss. 2014, 172, 413–420. 163. Kashiwagi, T.; Elsler, B.; Waldvogel, S. R.; Fuchigami, T.; Atobe, M. J. Electrochem. Soc. 2013, 160, G3058–G3061. 164. Elsler, B.; Schollmeyer, D.; Dyballa, K. M.; Franke, R.; Waldvogel, S. R. Angew. Chem., Int. Ed. 2014, 53, 5210–5213. 165. Wellmann, J.; Steckhan, E. Chem. Ber. 1977, 110, 3561–3571. 166. Sasaki, K.; Kunai, A.; Harada, J.; Nakabori, S. Electrochim. Acta 1983, 28, 671–674. 167. Ohnishi, R.; Aramata, A. J. Chem. Soc., Chem. Commun. 1986, 1630–1631. 168. Fleszar, B.; Sobkowiak, A. Electrochim. Acta 1983, 28, 1315–1318. 169. Lee, B.; Naito, H.; Hibino, T. Angew. Chem., Int. Ed. 2012, 51, 440–444. 170. Lee, B.; Naito, H.; Nagao, M.; Hibino, T. Angew. Chem., Int. Ed. 2012, 51, 6961–6965. 171. Polat, K.; Aksu, M. L.; Pekel, A. T. J. Appl. Electrochem. 2002, 32, 217–223. 172. Kariv-Miller, E.; Swenson, K. E.; Zemach, D. J. Org. Chem. 1983, 48, 4210–4214. 173. Chaussard, J.; Combellas, C.; Thiebault, A. Tetrahedron Lett. 1987, 28, 1173–1174. 174. (a) Benkeser, R. A.; Kaiser, E. M. J. Am. Chem. Soc. 1963, 85, 2858–2859. (b) Benkeser, R. A.; Kaiser, E. M.; Lambert, R. F. J. Am. Chem. Soc. 1964, 86, 5272–5276. 175. Sternberg, H. W.; Markby, R. E.; Wender, I.; Mohilner, D. M. J. Electrochem. Soc. 1966, 113, 1060–1062. 176. Sternberg, H. W.; Markby, R. E.; Wender, I.; Mohilner, D. M. J. Am. Chem. Soc. 1969, 91, 4191–4194. 177. Misono, A.; Osa, T.; Yamagishi, T.; Kodama, T. J. Electrochem. Soc. 1968, 115, 266–267.

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178. Ishifune, M.; Yamashita, H.; Kera, Y.; Yamashita, N.; Hirata, K.; Murase, H.; Kashimura, S. Electrochim. Acta 2003, 48, 2405–2409. 179. Tomilov, A. P.; Chernykh, I. N.; Zabusova, S. E.; Sigacheva, V. L.; Alpatova, N. M. Sov. Electrochem. 1987, 10, 1248–1255. 180. Kariv-Miller, E.; Swenson, K. E.; Lehman, G. K.; Andruzzi, R. J. Org. Chem. 1985, 50, 556–560. 181. Tountian, D.; Brisach-Wittmeyer, A.; Nkeng, P.; Poillerat, G.; Ménard, H. J. Appl. Electrochem. 2009, 39, 411–419. 182. Misra, R. A.; Sharma, B. L. Electrochim. Acta 1979, 24, 727–728. 183. Amouzegar, K.; Savadogo, O. Electrochim. Acta 1994, 39, 557–559. 184. Li, Z.; Garedew, M.; Lam, C. H.; Jackson, J. E.; Miller, D. J.; Saffron, C. M. Green Chem. 2012, 14, 2540–2549. 185. Misra, R. A.; Yadav, A. K. Bull. Chem. Soc. Jpn. 1982, 55, 347–348. 186. Shono, T.; Matsumura, Y.; Tsubata, K.; Sugihara, Y. J. Org. Chem. 1979, 44, 4508–4511. 187. Akbulut, U.; Toppare, L.; Utley, J. H. P. J. Chem. Soc., Perkin Trans. 2 1982, 391–394. 188. Mairanowski, W. G.; Wolkowa, O. I.; Obolnikowa, E. A.; Samochwalow, G. I. Dokl. Akad. Nauk SSSR 1971, 199, 829–831. 189. Ahmed, S.; Khan, A. Y.; Qureshi, R.; Subhani, M. S. Russ. J. Electrochem. 2007, 43, 811–819. 190. Guin, P. S.; Das, S.; Mandal, P. C. Int. J. Electrochem. 2011, 1–22. 191. Newton, L. A. A.; Cowham, E.; Sharp, D.; Leslie, R.; Davis, J. New J. Chem. 2010, 34, 395–397. 192. Abul-Hajj, Y. J.; Tabakovic, K.; Tabakovic, I. J. Am. Chem. Soc. 1995, 117, 6144–6145. 193. Wolf, R. E.; Cooper, S. R. J. Am. Chem. Soc. 1984, 106, 4646–4647. 194. Comninellis, C.; Plattner, E. J. Appl. Electrochem. 1985, 15, 771–773. 195. Mandell, L.; Cooper, S. M.; Rubin, B.; Campana, C. F.; Day, R. A. J. Org. Chem. 1983, 48, 3132–3134.

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27

Sulfur-, Selenium-, and Tellurium-Containing Compounds Richard S. Glass

CONTENTS I. II. III.

Introduction ........................................................................................................................ 1035 Thiols .................................................................................................................................. 1035 Disulfides............................................................................................................................ 1037 A. Reduction .................................................................................................................... 1037 B. Oxidation .................................................................................................................... 1041 IV. Thioethers (Sulfides) and Thioketals.................................................................................. 1045 A. Oxidation .................................................................................................................... 1045 B. Reduction .................................................................................................................... 1050 V. 1,4-Dithiin and Thianthrene............................................................................................... 1052 VI. Tetrathiafulvalene and Analogues...................................................................................... 1055 VII. Thiophene and Analogues .................................................................................................. 1072 VIII. Polythio and Sulfur-Nitrogen Unsaturated Heterocycles ................................................... 1076 IX. Sulfonium Salts, Sulfoxides, Sulfones, Sulfinyl, and Sulfonyl Derivatives ....................... 1078 X. Selenium and Tellurium Compounds ................................................................................. 1080 References .................................................................................................................................... 1087

I. INTRODUCTION The electrochemistry of organosulfur, selenium, and tellurium compounds is reviewed with particular emphasis on studies published from 2000 to the time of writing. There have been recent reviews of organosulfur electrochemistry [1,2], and this review brings them up to date and also includes the less studied electrochemistry of organoselenium and tellurium compounds. The reviewed work draws from many diverse areas, including organic synthesis, reaction mechanisms, and materials science. Therefore, in addition to summarizing the electrochemical results, insight into the significance of the electrochemical results to the relevant areas will be briefly presented.

II.

THIOLS

The oxidation of thiols depends on pH owing to the acidity of the thiol [1,2]. On increasing the pH, oxidation peak potentials become more negative until the thiolate is completely formed after which the oxidation potential is constant. Traditionally, mercury electrodes have been used for this reversible oxidation to the corresponding disulfides. This redox process occurs via the intermediacy of mercury thiolates [2]. Use of solid electrodes such as Pt, Au, or carbon for this process suffers from slow heterogeneous electron-transfer rates as well as adsorption (although a Pt electrode was recently used for cv studies of 2-pyrimidinethiols in ethanenitrile in which irreversible oxidation was reported to give 1035

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1036

Organic Electrochemistry

the corresponding disulfides and only on repetitive scans was adsorption identified) [3]. Consequently, several strategies have been developed to overcome these problems. These developments have been particularly motivated by interest in detecting biologically important thiols such as l-cysteine. Boron-doped diamond electrodes show negligible adsorption and favorable electron-transfer kinetics suitable for electrooxidation of thiols [4–6], and even proteins containing cysteine residues have been studied [7,8]. However, other residues in proteins are oxidized in addition to cysteine. Although carbon paste and glassy carbon electrodes typically suffer from the drawbacks outlined earlier, use of mesoporous carbon-modified glassy carbon electrodes permits detection of l-cysteine electrochemically [9] as do carbon nanotube electrodes [10]. Although use of platinum electrodes is also problematic, as mentioned earlier, platinum nanoparticles supported on nickel–cobalt nanofilms show high sensitivity and long-term stability for the determination of cysteine by anodic oxidation [11]. Homogenous redox mediators such as ferricyanide [12] and ferrocenyl trimethylammonium [13] have been used for l-cysteine oxidation using boron-doped diamond or glassy carbon electrodes, respectively. The redox catalyst has also been incorporated into the electrode for cysteine oxidation: quinizarine (1,4-dihydroxy-9,10-anthraquinone) adsorbed on glassy carbon [14], ruthenium (IV) oxide-modified carbon paste [15], or Nafion-lead oxide-manganese oxide [16] electrodes. 2-Aminoethanethiol [17] and glutathione [18] have been electrochemically oxidized using cobalt phthalocyanine adsorbed on a graphite electrode [17]. Similarly, 2-aminoethanethiol has been anodically oxidized using electropolymerized cobalt porphyrin and phthalocyanine-based films deposited on vitreous carbon electrodes [19]. An alternative approach has been developed for the electrochemical determination of thiols not involving oxidation of the thiol to the corresponding disulfide. Here N,N-dimethyl-1,4phenylenediamine in solution is electrochemically oxidized to the corresponding 1,4-diimine. The thiol then adds conjugatively to the 1,4-diimine to give an adduct, which is again electro-oxidized to a sulfur substituted 1,4-diimine as shown in Scheme 27.1 [20]. A monolayer of thiophenolate has been grafted onto glassy carbon and this surface-confined species can apparently be reversibly oxidized to the corresponding disulfide despite steric constraints [21]. Electrochemical oxidation of thiols to disulfides involves the intermediacy of thiyl radicals. Such an intermediate from 1,2,4-triazole-3-thiols has been trapped with dimethyl or diethyl acetylenedicarboxylate in an electrosynthesis of sulfur heterocycles as illustrated in Scheme 27.2  [22]. NH

NH2 –2e–, –H+

NH2

NH –2e–, –H+

RSH

+2e–, +H+

+2e–, +H+

SR NMe2 +

NMe2

SR

NMe2

NMe2 +

SCHEME 27.1 Electrochemical thiol determination with N,N-dimethyl-1,4-phenylenediamine. H CO2R΄

O

R

H N

S

H N

R

NH N R = Me, Ph

N

SH N

H N

R

S

–e– –H+

R΄O2CC N

N

R CCO2R΄

R΄= Me, Et

SCHEME 27.2 Trapping of a thiol radical with dialkyl acetylenedicarboxylates.

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N

N N 86–90%

S

1037

Sulfur-, Selenium-, and Tellurium-Containing Compounds

Anodic oxidation of C6S82−, 1a, and C6S6O22−, 1b, occurs in two reversible one-electron transfer steps as shown by cyclic voltammetric studies, to the corresponding disulfide 2a and b, respectively [23] S S

S

S

S

X

S– S

S

S

S

S –S

S X

S

S

S

X

S

S

S–

X 1a, X = S b, X = O

2

S

2a, X = S b, X = O

3

The one-electron oxidation product from 1a is stable in solution in equilibrium with disulfide dimer 3 [24]. While removal of an electron from an RS− moiety in 1a formally gives RS•, it has been reasonably suggested that the RS• moiety bonds intramolecularly with the RS− moiety forming the disulfide anion radical of 2a [24].

III.

DISULFIDES

A.

REDUCTION

As pointed out in Section II, electrochemical oxidation of thiols produces disulfides that may be electrochemically reduced back to thiols/thiolates and such electro reductions have been reviewed [25]. Owing to its commercial interest, the electrochemical reduction of l-cystine to l-cysteine has attracted particular attention. The key chemical issue is that this reduction must be run in acid because of the low solubility of these amino acids in aqueous solution of pH > 2. However, at low pH, reduction of protons to produce H2 becomes a competitive reaction. Consequently, electrodes at which there is a high overpotential for H2 production have been utilized. Early work has been reviewed [26], and there have been recent studies using lead [27], mercury [28], and dispersed lead– carbon black electrodes [29]. The thiolates produced by electrochemical reduction have been used as nucleophiles in synthetically useful reactions [2]. A recent example utilizes the thiolate produced by reduction of disulfide 4 to react with thio-glycosyl bromide 5a to produce 5b in 40–70% yield depending on reaction conditions [30]. This electrosynthesis is performed in an undivided cell using zinc as a sacrificial anode. AcO AcO

O NC

C

S

S X AcO

2 4

5a, X = Br b, X = S

O C

CN

Electrochemical reduction of disulfides played a key role in the construction of a monolayer of thiophenolate grafted onto glassy carbon whose electrooxidation was outlined in Section II [21]. Here, selective electrochemical reductions of the diazonium, not disulfide, moiety in diazonium salt 6 generated the corresponding radical that covalently bonded to the glassy carbon electrode surface as outlined in Scheme 27.3.

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1038

Organic Electrochemistry SSAr +N

S

2

X

S

e–

S

X

S 7

6, X = H, Cl

SSAr

n(2e–) S–

SCHEME 27.3 Electrochemical grafting of disulfides by diazonium salt reduction. RSSR + e– RS + e–

RS + RS– RS– 2RS–

Overall: RSSR + 2e–

SCHEME 27.4

Mechanism for electrochemical reduction of disulfides.

Typically such electrografting results in intertwining multilayers attached to the surface. However, here the more accessible aromatic ring remote from the surface preferentially reacts further with aryl radicals, for steric reasons, thereby protecting the inner ring yielding species such as 7. Electroreduction of the disulfides removes these superfluous moieties, resulting in a monolayer of phenylthiolate. This methodology has also been applied to functionalizing carbon nanotubes [31]. Unfortunately, further studies have revealed that there is not solely a homogenous monolayer of phenylthiolates because there is apparently some carbon–carbon and carbon–sulfur coupling leading to biphenyl dithiol- and phenylthiobenzenethiol moieties attached to the surface [32]. Comprehensive studies on the mechanism of electron transfer to disulfides have provided a treasure trove of understanding about the details of this process [33,34]. Addition of an electron to a disulfide results in a dissociative reduction in which the disulfide bond is cleaved to form a thiyl radical and thiolate. Addition of another electron results in reduction of the thiyl radical to thiolate. Thus, two-electron reduction of the disulfide produces two thiolates as shown in Scheme 27.4. The key issue concerns the details of the first step: addition of an electron to the disulfide. Does it occur stepwise, that is, via the formation of an anion radical or is the addition of an electron concerted with fragmentation? [35]. To address this question, a variety of techniques including voltammetry with heterogeneous electron transfer and homogeneous electron transfer via electrogenerated anion radicals, convolution voltammetry, and computational methods were employed [33,34]. Thus, a series of diaryl disulfides 8, X = NH2, MeO, H, F, Cl, CO2Et, CN, NO2 were studied [36–38]. N

X

ArCH2S

S

2

R

N

S

2 8

9

10a, R = SH b, R = Me

O

S

O

2 S 11

n

The heterogeneous and homogeneous electron transfer parameters correlated with each other, indicating that the reaction mechanism is the same for both reductions. Typically the reorganization energy ∆GO≠ for stepwise reductions is modest (2.3–3 kcal/mol) and is determined by the solvent ≠ ≠ . This is due to typireorganization energy ∆GO,s rather than the inner reorganization energy ∆GO,i ≠ cally low values for ∆GO,i because addition of an electron to an aromatic ring results in minimal

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1039

Sulfur-, Selenium-, and Tellurium-Containing Compounds

change in geometry, for example, bond lengths. On the other hand, if the reaction is concerted then ≠ ∆GO≠ is larger owing to an increased ∆GO,i in which the bond dissociation energy of the breaking ≠ bond contributes substantially. For most diaryl disulfides 8, ∆GO,i is high, suggesting concerted dissociative reduction. However, computations and other experimental evidence implicate the inter≠ mediacy of disulfide radical anions and a stepwise mechanism. Consequently, the large ∆GO,i is ascribed to the formation of a “loose radical anion” with a substantially elongated S–S bond in the σ∗ (2c, 3e bonded) anion radical [39]. These results bridge the previous work on stepwise and concerted limiting mechanisms for dissociative electron transfer and deepen our understanding of these ≠ processes. For diaryl disulfide 8, X = NO2, ∆GO,i is reduced—consistent with a conventional stepwise mechanism. Calculations suggest that, in this case, a conventional π∗ radical anion is formed that undergoes intramolecular electron transfer to the S-S moiety providing a σ∗ anion radical that then cleaves. Electrochemical studies on dibenzyl disulfides 9 also show stepwise dissociative electron transfer via a loose disulfide anion radical [40]. Reductions of disulfides have also been of interest for potential use in rechargeable lithium batteries [41–43] and organic dye-sensitized solar cells. Here, a particular concern is the sluggish kinetics of disulfide reduction at conventional electrodes, for example, glassy carbon. Redox reactions of disulfide 10a are dramatically accelerated at a glassy carbon electrode modified with a poly(3,4-ethylenedioxythiophene), PEDOT, 11 film than the unmodified electrode [44,45]. Similarly for reduction of 10b on a PEDOT film rather than on a platinized glass electrode [46]. Reduction of disulfides on gold electrodes is complicated by the formation of gold thiolates. This has recently been illustrated by comparing the reduction of diphenyl disulfide on a glassy carbon electrode with that on a gold electrode [47]. Cyclic voltammetric studies of di- and trisulfides 12a and 12b, respectively, have been reported on gold and gold coated with a single-atom monolayer of β-cyclodextrin thiol electrodes [48]. S Ph3CSXR

S–

S

l

l

l

l –

S

12a, X = 2 b, X = 3

13

14

The reduction of a disulfide moiety mediated by duplex DNA has been reported [49]. A monolayer of DNA, modified to bear a disulfide linkage, was assembled on pyrolytic graphite anchored by an appended pyrene moiety. Square wave voltammetric studies revealed an irreversible reduction centered at −160 mV and a reversible reduction at −290 mV versus NHE. The irreversible, but not the reversible, reduction is suppressed in acid. The reversible process is suggested to be a concerted 2e−, 2H+ reduction mediated by the DNA because a mismatch in the DNA duplex in the DNA closer to the electrode surface than the disulfide link substantially lowers the peak current but a mismatch after the disulfide does not. S

S–

S

H2N

NH2

H2N

NH2 S–

S

–S

–S

S 15

16

1,2-Dithiins represent disulfides in which reduction results in a dithiolate in which the two thiolate moieties are in the same molecule facilitating reoxidation to disulfide. Furthermore, the preferred geometry of the two species, disulfide and dithiolate, are different, and this has attracted interest

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1040

Organic Electrochemistry

for designing molecular memory devices and electrochemical switches. It also raises intriguing issues concerning the mechanism of electroreduction: specifically the timing of electron transfers and geometric changes. 1,2-Dithiin 13 has a dihedral angle of 34° about the S–S bond and an S–S bond length of 2.055 Å [50,51]. It undergoes this two electron reduction to dithiolate 14 in which the dihedral angle between the two aromatic rings is 130° and the S,S distance is calculated to be 5.35 Å. Clearly there is a dramatic change in geometry on reduction. Cyclic voltammetric studies on 1,2-dithiin 2b in DMF show two reversible reductions at −0.07 and −0.21 V versus SSCE [23]. The structures of the reduction products and their reoxidation were discussed in Section II. Irreversible electrochemical reduction of 15 in CH2Cl2 was reported [52] and presumed, but not proven, to produce 16. The interesting reduction of 1,2-dithiin fused to the 4,4′-bipyridinium system of methyl viologen has been reported. Four reversible one-electron waves are reported corresponding to overall two-electron reduction of the 4,4′-bipyridinium and two-electron reduction of the disulfide moieties [53]. The electrochemical reduction of sulfenyl chlorides and sulfenyl thiocyanates is discussed here because of their relevance to disulfide reduction. Electroreduction of arenesulfenyl chlorides 17 occurs with dissociative reduction to give the corresponding disulfides. O2N X

SCl Y

17a, X = Me, Y = H b, X = Y = H c, X = Cl, Y = H d, X = NO2, Y = H e, X = H, Y = NO2 f, X = Y = NO2

X

CH2SCN Y 18a, X = Me, Y = H b, X = Y = H c, X = MeO, Y = H d, X = Cl, Y = H e, X = F, Y = H f, X = CN, Y = H g, X = NO2, Y = H h, X = M, Y = NO2

SCN O2N

X

19

SCN

20a, X = Me b, X = OMe

Cyclic voltammetric studies show an irreversible one-electron reduction and similar voltammograms for 17a–d, on the one hand, and 17e and f, on the other hand, but the peak potential for 17e is 420 mV more negative than that for the isomeric 17d [54,55]. DFT computational studies suggest that electron transfer to 17a–d is concerted with S–Cl bond cleavage to yield a radical/anion pair (“sticky” dissociative electron transfer). Whereas electron transfer to 17e and f involves the formation of π∗ radical anions and subsequent bond cleavage, that is, a stepwise dissociative mechanism. Sulfenyl chlorides 17e and f form π∗ radical anions owing to S⋯O intramolecular interaction involving the o-nitro group that can occur in 17e and f but is geometrically precluded in p-nitro substituted 17d. Cyclic voltammetric studies of benzyl thiocyanates have been reported [56,57]. Compounds 18a–e undergo a one-electron irreversible reduction to produce the corresponding thiyl radicals that are reduced at a lower potential to thiolates. The thiolates, in turn, react with starting thiocyanate to form the corresponding dibenzyl disulfides that are more difficult to reduce than the corresponding thiocyanates. The addition of an electron to the benzyl thiocyanates is suggested to occur concertedly with S–CN bond cleavage, although CH2–S bond cleavage is also observed. Thiocyanate 18f behaves similarly except that only CH2–S not the S–CN bond cleaves. Compounds 18g and h undergo irreversible one-electron reduction, but their electrochemical behavior suggests stepwise electron transfer. DFT computational studies suggest formation of an aromatic π∗ radical anion in these cases followed by CH2–S cleavage. Product studies reveal the corresponding ArCH2CH2Ar resulting from CH2–S cleavage of the radical anion. Addition of phenol results in the formation of ArCH3 instead of ArCH2CH2Ar. Consequently, it is suggested that the ArCH2 formed by CH2–S cleavage of the thiocyanate radical anion does not form ArCH2CH2Ar by dimerizing.

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1041

Sulfur-, Selenium-, and Tellurium-Containing Compounds +e– ArSCN

ArS +e– –

+e–

CN

ArS–

ArSSAr –CN

ArSCN

SCHEME 27.5 Autocatalytic reduction of aryl thiocyanates.

Rather the benzyl radical is reduced to the corresponding anion that attacks the benzyl thiocyanate displacing −SCN and forming ArCH2CH2Ar. Voltammetric studies and convolution analysis of the electrochemical reduction of aryl thiocyanate 19 show a stepwise mechanism via a π∗ radical anion. Investigations with 20a and b suggest electron transfer is borderline between a stepwise and a mechanism concerted with S–CN cleavage [58]. However, an autocatalytic process was also found. In these cases, the reduction potential of the product diaryl disulfide is less negative than that for the corresponding aryl thiocyanates. Consequently, the arylthiolates, produced by aryl disulfide reduction, chemically react with the starting arylthiocyanates to produce diaryl disulfide, resulting in a catalytic cycle as shown in Scheme 27.5. The result of this autocatalytic mechanism is a lowering of the potential for reducing aryl thiocyanates.

B. OXIDATION Electrochemical oxidation of disulfides has been reviewed recently [1,2]. Consequently, the emphasis here is on results published subsequent to the prior reviews. Disulfides may undergo one- and two electron oxidation [59,60], depending on the solvent and nature of the attached groups: alkyl versus aryl, under anhydrous conditions and in the absence of oxygen donors, to afford cationic species. The structures of these cationic species have not been established definitively but they function as RS+ donors [61]. Although, recently, detection of diaryl disulfide radical cations generated electrochemically has been achieved in CH2Cl2 with tetrabutylammonium tetrakis(pentafluorophenyl)borate as supporting electrolyte [62]. The mechanistic complexities of electrochemical oxidation of disulfides are well illustrated in detailed studies on anodic oxidation of disulfides 21–23 [63–65]. O O Me N O

(CH2)n OCCH2 S

RN N

N Me

N

O (CH2)nS O O

21, n = 2, 3, 4, 6

2

N

O N

S

MeN

S O S

N

(CH2)n OCCH2 O 22, R = Me, Ph; n = 2, 3

(CH2)3S(CH2)2

O

O

(CH2)3S(CH2)2 23

While mechanistic aspects of disulfide oxidation are complex, the synthetic use of the oxidized species as RS+ donors, a potent electrophile, has been very fruitful. Thus, controlled potential electrolysis of dimethyl disulfide or diphenyl disulfide in dichloromethane followed by reaction with phenols or aromatic ethers results in electrophilic aromatic substitution as illustrated in Equation 27.1 [61,66]. Such electrochemical methylthiation reactions occur in better yield using liquid SO2 instead of dichloromethane as solvent, p-substitution is favored over o- and weakly activated arenes react as well [67].

© 2016 by Taylor & Francis Group, LLC

1042

Organic Electrochemistry OR΄

OR΄ R˝

R˝

R˝

R˝

(27.1)

+ RS+-donor

SR

Controlled potential electrolysis of disulfides until the charge corresponding to two-electron oxidation is passed followed by the addition of a thiol generates unsymmetrical disulfides [68]. Reaction of electrochemically generated RS+ donors with alkenes provides thiiranium ions that can be trapped with a variety of nucleophiles both inter- and intramolecularly, resulting in antiaddition of sulfur and nucleophile [1,2,69]. Further advances in the synthetic application of these reactions have been made by conducting the controlled potential electrolysis of diaryl disulfides in dichloromethane containing Bu4NBF4 as supporting electrolyte with alkenes at low temperature. Reaction with alkenes at low temperature followed by addition of a nucleophile such as water or methanol gave addition products with regio- and stereochemical control as exemplified in Equation 27.2 [70] via the corresponding thiiranium salt. Apparently the dialkyl disulfide formed after transfer of ArS+ can react with the intermediary thiiranium ion because addition of Et3N, Me2C=C(OTMS)OMe or 3-trimethylsilylcyclohexene results in the formation of 1,2-arylthio addition products [70]. OH

H2O

(p–FC6H4S)3+ +

(27.2)

SC6H4pF

A reasonable mechanism for this reaction is shown in Scheme 27.6. Proposed intermediate 24 is an ArS+ donor. Consequently, (ArS)3+ can catalyze the addition of ArSSAr to alkenes. Indeed, 0.2 equivalents of (ArS)3+ catalyze the stereospecific antiaddition of ArSSAr to cis- and trans-1-phenyl propene in good yield as shown in Scheme 27.7 [71]. Alternatively, these catalytic reactions can be carried out by passing a catalytic amount of current in a solution containing ArSSAr and alkene. These disulfide additions work well with B(C6F5)4− but not BF4− as the counterion. Addition of ArSSAr to alkynes also occurs stereoselectively as ArS Ar S+ ArSSAr

+ (ArS)3+

24

ArSSAr +

ArS SAr

SCHEME 27.6 Mechanism for the formation of 1,2-di(arylthio) addition products. Ph

+

ArSSAr

Me

Ph

Me

+

ArSSAr

(ArS)3+ 0.2 equiv

(ArS)3+ 0.2 equiv

Ph

SAr

ArS

Me

Ph

Me

ArS

SAr

SCHEME 27.7 (ArS)3+ catalysis of stereospecific ArSSAr addition to alkenes.

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1043

Sulfur-, Selenium-, and Tellurium-Containing Compounds ArS +

+

(ArS)3

ArSSAr

CMe2SAr

0.2 equiv 77% ArS +

CHPh

+

(ArS)3

ArSSAr

CH(Ph)SAr

0.2 equiv 58% ArS

Ph

Ph

+

Ph

+

(ArS)3

ArSSAr

CH(Ph)SAr

0.2 equiv

SCHEME 27.8 Intramolecular cyclization of thiiranium ions with alkenes.

illustrated in Equation 27.3 [70]. If the reaction of electrochemically generated (ArS)3+ with alkenes or alkynes is done at 0°C rather than −78°C, BF4− can act as a fluoride donor, resulting in thiofluorination [72] as illustrated in Equation 27.4: Ph

PhC CR + ArSSAr R = Ph, Me, H

SAr

ArS

(27.3)

R F

RCH

CH2 + (ArS)3+BF4–

(27.4)

RCHCH2SAr 52–99%

Thiiranium ions produced under these conditions can react intramolecularly with alkenes as nucleophiles. Thus, the reactions shown in Scheme 27.8 proceed regioselectively [71]. Reactions of electrochemically generated (ArS)3+ in which ArS+ is donated have been extended to enolizable ketones, enol acetates, ketene silylacetals, and allyl silanes [72]. These electrochemically produced ArS+ donor species are not only potent carbon electrophiles but also thiophilic electrophiles are well. Thus, monothioacetals 25 react rapidly with electrogenerated (ArS)3+ at −78°C to afford alkoxycarbenium ions 26. These species in turn react with carbon nucleophiles, such as allylsilanes, silyl enol ethers, silyl ketene acetals, and enol acetates to form carbon–carbon bonds. An example is shown in Equation 27.5 [73]. In addition, monothioacetals 27–29 have been also used advantageously in this reaction. OMe +

(ArS)3

+ RCHSPh R

SiMe3 RCH

+ OMe

C18H17, Ph 25

OMe RCH – CH2CH

CH2

(27.5)

69–98% 26

Styrene also reacts with alkoxycarbenium ions 26 to ultimately produce thiochromans in a series of steps involving reaction with the diaryl disulfide in solution [74].

© 2016 by Taylor & Francis Group, LLC

1044

Organic Electrochemistry OMe SPh O 27

O

SPh

28

SPh

29

Intramolecular cyclization involving alkenes could be effected as well. Thus, synthesis of tetrahydropyrans 30 could be accomplished as shown in Equation 27.6 [75]: (ArS)3+

O R R΄

ArS

0.2 equiv –78°

(27.6)

O R΄ SAr

R

30 62–88%

R = R΄= H R = C7H15; R΄= Me R = C7H15; R΄= H R = c-Hex; R΄= H

R = Ph; R΄= H R = PhCH2; R΄= H R = MeOCH2CH2; R΄= H

Glycosyl thioethers are thioacetals and glycosyl cations can be formed by treatment of glycosyl thioethers with (ArS)3+ [76]. However, glycosyl cations are less stable than the cations obtained from the monothioacetals described earlier. Consequently, a flow microreactor system was developed to rapidly prepare these cations and couple them with sugar alcohols to synthesize disaccharides [77]. A perspective on the preparation of (ArS)3+ by electrolysis of ArSSAr at low temperature and its synthetic uses has recently been published [78].

R

R΄ S

S

31a, R = CH2 = CHC CC C–; R΄= MeC C– b, R = R΄ = H; c, R = R΄= Ph; d, R = R' = CH2OH; e, R = R΄= Me f, R = R΄= t-Bu; g, R = R΄ = i-Pr

S

S 32

Electrochemical studies on 1, 2-dithiin 31 in CH 2Cl2 and CH3CN report reversible one-electron oxidation followed by an irreversible oxidation [79] with peak potentials for 31a–d the range of 0.80–1.04 V for the first oxidation and 1.13–1.40 V for the second in CH 2Cl2 versus Ag/0.1 M Ag+ and 0.58–0.96 and 0.99–1.26 V, respectively, for 31a–g in CH3CN. Interestingly, neutral 1,2-dithiins 31 are calculated to be nonplanar while the corresponding radical cations are calculated to be planar or near planar. Consequently, the first oxidation occurs by an EC mechanism for which electrochemical parameters show the electron transfer and geometry change to be sequential. 1,2-Dithiin 32 undergoes two reversible one-electron oxidations in CH 2Cl2 at −78°C with E1/2 + 0.18 and + 0.72 versus Fc/Fc+ (at room temperature the second oxidation is irreversible) [80]. The lowered first oxidation potential for 32 compared with 3,6-dialkyldithiins (oxidation peak potentials are estimated to be +0.46 to +0.53 V vs. Fc/Fc+ from potentials in CH3CN vs. Ag/0.1 M Ag+ reported in Reference 79) are ascribed to inductive effects and C–C σ–π conjugation. Cyclic voltammetric studies on a dithiete fused to a sterically hindered o-quinone show two reversible 1e-reductions centered on the quinone moiety and one quasi-reversible oxidation [81].

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1045

Sulfur-, Selenium-, and Tellurium-Containing Compounds

IV. THIOETHERS (SULFIDES) AND THIOKETALS A. OXIDATION Electrochemical oxidation of thioethers and thioketals has been reviewed previously [1,2]. Anodic oxidation of thioethers under anhydrous conditions and typical cyclic voltammetric conditions involves an EC mechanism. Usually peak potentials for such oxidations are reported. However, using second harmonic ac voltammetry, formal (thermodynamic E1) potentials for one-electron oxidation of a variety of thioethers in acetonitrile were reported [82]. The chemical step following the formation of sulfur radical cations by one-electron oxidation is nucleophilic attack, loss of α-proton, or bond cleavage (C–S, α C–M, decarboxylation) [83]. Nevertheless, sulfur radical cations can be stabilized by their propensity to form 2c-, 3e-bonds. Such a propensity not only stabilizes such species as the one-electron oxidation product, such as that formed from 1,5-dithiocane 33, resulting in electrochemically reversible oxidation, but substantially less anodic (by ca. 1 V) oxidation potentials and potential inversion [84].

S

COX

S MeS

COX

S +

MeS Me

33

b, X = N(CH2)4

X

36a, X = NH2

35a, X = NH2

34a, X = NH2

C

O

b, X = N(CH2)4

b, X = N(CH2)4

Removal of a second electron is less anodic than the first by about 0.15 V. This result is ascribed to 2c, 3e-S, S-bonding in the radical cation of 33, illustrated in Scheme 27.9, in which an antibonding σ* electron is removed to afford the corresponding dication [84,85]. Two-center, three-electron bonds between sulfur and other heteroatoms are also known [39]. Consequently, juxtaposing moieties with such atoms close to thioethers may render their anodic peak potentials less positive. Thus, endo-amide 34a, in which the sulfur is held close to the amide group, oxidizes irreversibly at a peak potential 330 mV less positive than exo-amide 35a, in which these moieties are precluded from forming a 2c, 3e S, O-bond [86]. Pulse radiolysis studies confirm the formation of transient 36a on one-electron oxidation of 34a. That 37 and 38, whose geometry permits S,O interaction but precludes S,N-interaction, show similar facilitated oxidation potentials as 34a provides additional evidence that 2c, 3e S,O-bonds, rather than 2c, 3e S,N-bonds, are formed on one-electron oxidation of endo-amide 34a. H N

H N

O

O MeS

O

MeS

N H

37

O

CONH2 SMe

N H 38

39

σ* p

p σ

SCHEME 27.9 2c, 3e-Bonding.

© 2016 by Taylor & Francis Group, LLC

1046

Organic Electrochemistry

In the presence of trace amounts of bromide ion, 34a, but not 35a, shows electrochemical redox catalysis. Amide 39, a regioisomer of 34a in which 2c, 3e S,O-bond formation results in a fivemembered rather than a six-membered ring, also undergoes facilitated oxidation. Again pulse radiolysis studies confirmed the formation of a transient with a 2c, 3e S,O-bond on one-electron oxidation of 39, which is even longer lived than 36a. Tertiary amides, as in 34b, are more electron-rich than primary amides, as in 34a, therefore, one might anticipate a less positive oxidation potential for 34b than for 34a. Indeed, the irreversible oxidation potential for 34b is 330 mV less anodic than that for 34a, 660 mV less positive than that for 35a, and 530 mV less positive than that for 35b [87]. Aromatic rings juxtaposed with thioethers as in 40 undergo irreversible anodic oxidation at less positive peak potentials than their isomers 41, in which through space interaction is precluded [88].

Ar SMe Ar

MeS

X

X

MeS

40

41

42

DFT calculations suggest that the lowered oxidation potentials are due to interaction of the sulfur lone pair and aromatic π MOs. That is, the HOMO from which the electron is removed in 40 is a combination of sulfur 3p and arene π orbitals. m-Terphenylthioethers 42 preferentially adopt a conformation that favors sulfur 3p-arene π interaction as shown [89]. This results in less anodic oxidation potentials [90]. An important factor contributing to the lowered electrochemical oxidation potentials for 40 has been reported [91] to be due to a new type of S∴π bonding in the corresponding radical cation. Interest in applying electrochemical oxidation of thioethers and the subsequent reactions of the sulfur radical cations to synthesis has continued, and recent developments will be reviewed. Electrochemical oxidation of thioethers followed by α-deprotonation gives an α-thioradical 43 that undergoes one-electron oxidation (at lower potential than the corresponding thioether) to yield thionium ion 44 as shown in Equation 27.7. An alternative mechanism for the formation of thionium ion 44 by electrochemical oxidation of sulfides is shown in Equation 27.8. RCH2SR΄

–e–

–H+

+

RCH2SR΄

–e–

RCHSR΄

+

RCH =SR

43

(27.7)

44

Thionium ions 44 are believed to be intermediates in the Pummerer reaction [92–96]. X –e– RCH2SR΄

X– –e–

+

RCH2SR΄

–HX RCH2SR΄ +

44

(27.8)

45

They are electrophilic and react both intra- and intermolecularly with a variety of nucleophiles. Electrochemical oxidation of thioether 46, X = H in the presence of methanol, acetate, or fluoride affords the corresponding derivative 46, X = OMe, OAc or F [97]. PhSCH(X)CF3 46

Here, the CF3 group promotes deprotonation. Similarly, other electron-withdrawing group esters [98,99], amides [98], ketones [98,99], cyano [98], acetylene [100], phosphonate [98,101,102], and sulfonyl [103] promote deprotonation and regioselectivity in anodic fluorination. The diastereoselectivity

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1047

Sulfur-, Selenium-, and Tellurium-Containing Compounds

of these reactions has been investigated [104,105] and a particularly impressive example in which only diastereomer 47 is detectable by 1H NMR spectroscopy is illustrated in Equation 27.9 [106]. –2e, –H+ S R2NSO2

O

DME 88% yield

O

F

S

Et4NF . 4HF R2NSO2

(27.9)

O O 47

Anodic α,α-difluorination has also been reported, for example, anodic fluorination as shown in Equation 27.10 [98]. –2e–, –H+ PhSCH2CO2Et

–2e, –H+ PhSCH (F)CO2Et

F–

48

PhSCF2CO2Et

F–

(27.10)

49

Occasionally anodic fluorination results in fluorination of the aromatic ring appended to sulfur as with oxazolyl [107] and thiazolyl [108]-2-sulfides rather than Pummerer reaction. Electrochemical α-fluorination of thioethers and anodic difluorination accompanied by desulfurization of dithioacetals has been included in a review of fluoro-Pummerer rearrangements [109]. A review focused on these reactions, as well as those carried out in ionic liquids, has also appeared [110]. In cases in which acidifying substituents are not attached to the carbon α to the sulfur, anodic α-fluorination can still be achieved. The suggested explanation is that the reaction course follows the mechanism shown in Equation 27.8 in which the electron-withdrawing fluorine in intermediate 45, X = F contributes to acidifying the α-proton [108,111–114]. Indeed, this explanation also accounts for the reported promotion of anodic methoxylation and acetoxylation of thiazolidines, 1,3-oxathiolanes and 1,3-dithiolanes [115]. Even more remarkable than anodic fluorination of unactivated thioethers is α-fluorination in systems in which competing C–S bond cleavage of the intermediary sulfur radical cation is promoted by substituents that stabilize the carbocation resulting from such C-S cleavage. For example, anodic oxidation of monothioacetals, such as phenyl-thioglucoside 50a, generates a radical cation which on C–S bond cleavage gives an oxygen stabilized carbocation, that is, oxonium ion [112]. OAc O

AcO AcO AcO

BnCONH

OAc X

AcO AcO

F AcO

SPh

S SAr

N O CO2PMB

50a, X = SPh b, X = F

51

52

Indeed, anodic fluorination with Et3N · 4HF produces fluoride 50b. However, anodic fluorination with Et3N·3HF in THF as solvent affords 51 as a mixture of anomers [112]. Thus, a number of factors have been found to affect the product formed and its yield. In the anodic methoxylation of cephem derivatives 52, the nature of the aryl substituent plays a key role in product selectivity [116]. As already mentioned, solvent [112,117,118] supporting electrolyte and fluoride source (Et3N · nHF, Et4NF · nHF) [112,119–124] are important as are use of mediators: Ar3N [125] or iodoarene incorporated into ionic liquid [126], ultrasonication [127], and use of additives such as the polyether PEG [128].

© 2016 by Taylor & Francis Group, LLC

1048

Organic Electrochemistry

To avoid the necessity of handing HF in anodic fluorination, an in situ method for generating it has been introduced [129]. The idea is that alkali metal fluorides will generate HF in acetonitrile solution with solid supporting sulfonic acids (cation exchange resins in the acid form). Furthermore, 2,6-lutidine is added so that it will react with HF to generate salts that will act as supporting electrolyte and fluoride source. Thus, anodic oxidation of ethyl α-phenyl-thioacetate 48 in acetonitrile solution containing KF, 2,6-lutidine and Amberlyst 15 Dry resin at reflux produces the corresponding α-fluoride 49 in 60% isolated yield. A comparable yield of the monofluorinated product was also obtained from PhSCH2CN. A novel approach to avoid the use of supporting electrolyte in the anodic methoxylation [130,131] and acetoxylation [132] of 46 has been reported. The idea is to use recyclable solid supported bases that form salts by acid–base reaction between the immobilized base and methanol or acetic acid in acetonitrile. The yields for such reactions of 46 with methanol and acetic acid are 89% and 81%, respectively. Such anodic methoxylation of PhSCHF2 has also been achieved in 86% yield [133]. The electrochemistry of hexakis(benzylthio)benzene 53 has been studied by cyclic voltammetry in CH2Cl2 [134]. The corresponding radical cation is formed on one-electron oxidation and characterized as a π-delocalized system by EPR spectroscopy. Analysis of fractional electrolysis (potentiometric titration) of this compound shows that it undergoes a second electron transfer with strong potential compression for the two one-electron transfer steps (E1° = 0.798 V, E2° = 0.821 V vs. Ag/0.01 M AgClO4 · H2O, 0.1 M Bu4NPF6, CH3CN from simulation) with ΔE° = 23 mV. As already pointed out, sulfur radical cations are known to undergo C–S bond cleavage to form carbocations particularly when the carbocation so formed is stabilized. Thus, electrochemical oxidation of thioethers followed by C–S bond cleavage has also been synthetically exploited. Electrolysis of benzylic sulfides in nitroalkane solutions in an undivided cell results in oxidation and C–S bond cleavage at the anode and α-nitrocarbanion formation at the cathode. Carbon–carbon bond formation then ensues as shown in Equation 27.11 (p-OMe or p-OH groups are required in the Ar moiety for successful reactions) [135]. electrolysis

→ ArCH(R)CH(R′′)NO2 ArCH(R)SR′ R ′′CH2 NO2 R′ = Me,Ph

(27.11)

Anodic oxidation of arylthioether derivatives of cholesterol and, even better 6β-3α,5α–cyclocholesterol, in the presence of partially protected monosaccharides provides glycoconjugates of cholesterol in moderate yields [136]. In this case, heterolysis of the C–S bond in the corresponding sulfur radical cation is favored by stabilization of the carbocation. Anodic oxidation of benzylic sulfides 54 results in C–S as well as C–Sn bond cleavage to generate o-quinodimethanes 55, which can be trapped with in situ dienophiles to give Diels–Alder adducts in excellent yields [137]. SCH2Ph PhCH2S

SCH2Ph

CH2SnBu3

CH2

PhCH2S

SCH2Ph

CH(R)SPh

CHR

ArCOX SCH2Ph 53

54

55

56a, X = SAr΄ b, X = OMe

Anodic oxidation of S-arylthiobenzoates 56a mediated by a triarylamine occurs with C–S bond cleavage to yield the corresponding methyl ester 56b in the presence of methanol [138] or benzoyl fluoride (from 56a, Ar = Ph) if Et4NF·3HF is used as the supporting electrolyte [139]. As described earlier, α-deprotonation and C–S bond cleavage are often competitive pathways and conditions favoring C–S bond cleavage over α-deprotonation have been reported [112,116,118,125].

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1049

Sulfur-, Selenium-, and Tellurium-Containing Compounds

Cleavage of the C–S bond in monothioacetals is favored by stabilization of the carbocation produced by the oxygen atom, although as reported earlier, conditions that favor α-deprotonation are known [112]. As previously reviewed, anodic oxidation of monothioacetal 57 in the presence of allyltrimethylsilane produces 58 via alkoxycarbenium ions resulting from C–S bond cleavage [1]. More recently, it has been shown that alkoxycarbenium ions are sufficiently stable that solutions of them can be electrosynthesized by anodic oxidation of monothioacetals in CH2Cl2 with Bu4NBF4 as the supporting electrolyte at −78°C. These electrophilic species then react with carbon nucleophiles (such as allylsilanes and silyl enol ethers) to form carbon–carbon bonds as illustrated in Equation 27.12 [140,141]. SiMe3 C8H17CH(OMe)SPh 57

Anodic ox Bu4NBF4 CH2Cl2 –78°

+

C8H17CH = OMe

C8H17CH(OMe)CH2CH=CH2

(27.12)

58

An alternative approach to this valuable methodology is to synthesize the alkoxycarbenium ion indirectly. Here, diaryl disulfides are electrooxidized to give an electrophilic sulfur species (see Section III.B) believed to be [ArS(ArSSAr)]+. Reaction of this species with monothioacetals affords the corresponding alkoxycarbenium ion [141]. Thioglycosides are monothioacetals and their electroxidation has attracted attention because C–S cleavage results in a glycosyl cation that on reaction with a sugar alcohol provides a disaccharide [1]. Efforts to improve the methodology have been reported [142], including a recent study on solvent and protecting group effects on the yields and stereochemistry at the anomeric center [143]. The oxidation peak potentials of the thioglycoside depend on the aryl substituent attached to sulfur [142,144]. This enables selective activation of glycosyl donors. A detailed electrochemical study of the formal (thermodynamic) potential of such glycosyl donors has been studied and found to correlate with peak potentials [145]. A review on the synthetic application of anodic oxidation of thioglycosides was recently published [146]. It would be advantageous to generate solutions of glycosyl cations by anodic oxidation of thioglycosides as was done with simple monothioacetals as outlined earlier. However, glycosyl cations are not very stable. Consequently, solutions of such cations cannot be electrosynthesized. For example, anodic oxidation under the conditions used for simple monothioethers, CH2Cl2, Bu4NBF4, −78°C, apparently provides the glycosyl fluoride. Use of Bu4NClO4 as the supporting electrolyte instead of Bu4NBF4 apparently gives the glycosyl perchlorate that reacts with sugar alcohols to yield disaccharides. Even better is to use Bu4NOTf as supporting electrolyte to provide solutions of glycosyl triflates [147,148]. It should be noted that it has also been reported that anodic oxidation of thioglycosides with 12.5 mol% sodium triflate as the supporting electrolyte in acetonitrile in the presence of sugar alcohols provides disaccharides [149]. Similarly, reaction of solutions of glycoside triflates with sugar alcohols affords disaccharides and even a pentasaccharide [147,148]. Although glycoside triflates are too reactive to be stored and must be used immediately, they can be converted to glycosyl sulfonium salts that can be stored and subsequently used in coupling reactions [150]. Electrocatalysis by Br− (or NBS or Br2) of the formation of nucleosides from protected thio-ribosides [151] and 1-arylthio-2,3dideoxy riboside [152] with silylated pyrimidines as shown in Equation 27.13 has been reported. O R΄˝ NH OTMS

ROCH2 O

ROCH2

R΄˝

SAr

N O

N + R΄

R˝ N

© 2016 by Taylor & Francis Group, LLC

OTMS

R΄

R˝

O

(27.13)

1050

Organic Electrochemistry

Anodic oxidation of dithioacetals and ketals has been reviewed previously [1,2]. More recently, intermolecular C–C bond formation has been achieved as shown in Equation 27.14 forming 59 in 80–90% yield [153]. OMe

OMe Anodic ox. LiClO4

R

MeO

R

MeO

(27.14)

CH3NO2 SiMe3

PhSCHSPh R = H, OMe

PhSCHCH2CH = CH2 59

Anodic oxidation of dithioacetal 60 in the presence of substituted alkenes also results in C–C bond formation.

O

OH CH(SPh)2

PhS 61

60

For example, anodic oxidation of 60 in the presence of tetramethylethylene gives 61 in 82% yield [153]. Anodic oxidation of dithioketals derived from diaryl ketones mediated by an iodobenzene ionic liquid [126], triarylamines [154], PhIClF [155], and Et3N·3HF as supporting electrolyte and F− donor as well as polymer supported iodobenzene and Cl− with Et3N·HF [156] produce the corresponding difluoride Ar2CF2. Oxidation of benzoyl ketene dithioacetals by a presumed ECEC mechanism has been reported to give mainly the product shown in Equation 27.15 [157] in 31% yield. Anodic ox PhCOCH=C(SMe)2

Et4NF . 4HF DME

PhCOC(F)2COSMe

(27.15)

B. REDUCTION Electrochemical reduction of aryl alkyl thioethers typically occurs with C–S bond cleavage in an overall two-electron process. For example, cathodic reduction of 62 proceeds as shown in Equation 27.16 [158]. The details for the dissociative electron transfer in such reductions have been investigated for Ph2CHSC6H4pOMe and found to be a stepwise process for the electron transfer with a large reorganization energy [159]. O(CH2)3SPh

O(CH2)CH3 + PhS–

OCH2OMe 62

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OCH2OMe 63

(27.16)

1051

Sulfur-, Selenium-, and Tellurium-Containing Compounds

As the structure of the sulfide is varied, stiff π* radical anions or loose σ* radical anions are formed [40]. Reduction of thioethers using mediators was found to be advantageous. For example, reduction of 62 mediated by naphthalene in THF at low temperature occurred at a potential 500 mV more positive than that for direct reduction [160]. Furthermore, the mediated reduction was a 1e − process providing PhSSPh in addition to 63. 4,4′-Di-t-butylbiphenyl was also an effective mediator for electro-reduction of sulfides [161]. Apparent electrochemical reduction of benzyl or tritylthiophenyl modified nucleotides and DNA in the presence of cobalt ions has been reported [162]. Cathodic reduction of aryl polythioethers has provided very interesting results. First, the reduction occurs at surprisingly low potentials and is reversible. For example, 1e− reduction of 64, Ar = Ph occurs at a formal potential of −1.56 V [163] −1.60 V [164] in DMF versus SCE. PhS

SPh Ar X Ar 66a, Ar = C6(SPh)5; X= – CH = CHC6H4CH = CH– b, Ar=C6(SPh)5; X = –Ch = N–N = CH– c, Ar=C6(SPh)5; X = –C C – C C–

SAr SAr

ArS

S

S SPh

PhS ArS

SAr

N

SAr

O O

O O

N

65

64

The reduction potentials of 64 with p-substituents in the Ar groups afforded a linear Hammett plot using σp parameters for these substituents [163]. Compound 65 has been studied as a redox sensor for K+ [165]. The idea is that the reduction of the aryl polythioether moiety will be affected by the binding of K+ to the cryptate moiety. Indeed, it was found that for 65 the binding constant for KPF6 was about 4000 M−1 and that there was a +170 mV shift for the reduction of this complex compared with uncomplexed 65. Electrochemical studies on linked perphenylthiobenzenes provided insight into the coupling of these centers. Thus, 66a and b showed 2e− reductions at −1.39 and −1.19 V, respectively, but 66c showed two 1e− reductions at −1.12 and −1.38 V in DMF versus SCE [166,167]. Therefore, the two Ar centers in 66a and b do not interact (class I in the Robin–Day classification system [168]), but addition of an electron to 66c results in an anion radical spread over the two Ar systems making it more difficult to add a second electron (borderline class II/class III system). This conclusion is validated by spectroelectrochemical studies in which the anion radical of 66c shows an intervalence charge-transfer band at 1310 nm in the near IR. This suggests the possibility of a molecular wire based on this structural motif and molecular rods 67 have been made and studied by cyclic voltammetry [169]. X Ar'S

X

SAr'

X X

X X

X Ar

X

X

X Ar

X

X Ar'S

SAr'

n

X

X

67a, Ar = C6(SPh)5; Ar' = Ph; n = 0 b, Ar = C6(SC6H4p–tBu)5; Ar' = C6H4p–tBu, n = 0 c, Ar = C6(SC6H4p–tBu)5; Ar' = C6H4p–tBu, n = 1

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X

X X

X X

68a, X = p –TolS b, X = iPrO n = 1-3 c, X = H

X

X

S

S; n

X

69a, X = SPh b, X = H

1052

Organic Electrochemistry

Compounds 67a and b show two reversible 1e− reductions at −1.12, −1.38 and −1.17, −1.44 V, respectively, in DMF versus Fc/Fc+ compared with 64, Ar = Ph or Ar = C6H4tBu, which show one reversible reduction at −1.56 and −1.7 V, respectively, under the same conditions. Extending the conjugated system to that in 67c results in a 1e− reduction at −0.98 V and a 2e− reduction at −1.19 V. Alternatively, arenepolythioethers with m-diacetylene bridges, in which the m-nature of the linkages prevents the possibility of full interaction, have been studied by cyclic voltammetry. In these cases, their multiple reversible reductions enable them to reversibly store several electrons, thereby acting as molecular batteries [170,171]. Cyclic voltammetric studies of perarylthiolated coronenes 68a and b also show two reversible 1e− reductions in DMF at potentials over 1 V less cathodic than coronene 68c [163,172] owing to stabilization by the sulfur substituents. Cyclic voltammetric studies of perphenylthiocorannulene 69a in acetonitrile show four reduction waves at −1.22, −1.62, −2.04, and −2.28 V versus Ag/AgNO3 [173]. Again the sulfur substituents lower the reduction potentials compared with the parent corannulene 69b (−2.23 and −2.84 V under comparable conditions). Here, FMO analysis provides further insight into the role of the sulfur substituents in rendering the first reduction potential less cathodic. The LUMO of 69a includes the sulfur atoms as well as the core π system, thereby lowering its energy and consequently the reduction potential.

V. 1,4-DITHIIN AND THIANTHRENE The electrochemistry of 1,4-dithiins and its derivatives has been reviewed previously [1,83,174]. More recently, the electrochemistry of substituted 1,4-dithiin 70 has been studied in dichloromethane by cyclic voltammetry [175,176]. X

S

X

X

S

X

S S 70

71a, X = H b, X = OMe

A reversible one-electron oxidation is followed by an irreversible oxidation [175]. However, at −78° both oxidations are reversible with E11/ 2 = +0.00 and E12/ 2 = +0.82 V versus Fc/Fc+[175]. The annulated bicyclo[2.2.2]octane framework stabilizes the oxidation products and the oxidation potentials are lowered owing to inductive and σ–π conjugation effects as suggested earlier for 1,2-dithiin derivative 32. Such effects by the bicyclo[2.2.2]octene moiety have been reviewed [177]. Electrochemical studies on thianthrene, 71a, and its derivatives have been reviewed previously [1,2,178]. Thianthrene undergoes two reversible, one-electron oxidations but the corresponding dication reacts with even trace amounts of water rendering the second oxidation of thianthrene irreversible. However, in dry solvents, the second oxidation is reversible [178–180]. In addition, zeolite modified electrodes show reversibility even in the presence of water owing to zeolite protection of the dication [181]. The PF6 − salt of thianthrene radical cation was electrosynthesized in CH2Cl2 with tetrabutylammonium hexafluorophosphate as the supporting electrolyte [182]. X-ray crystal structural analysis showed planar cations associated as dimers with two weak S,S bonds. Aggregation of this and related radical cations have been studied computationally [183]. 2,3,6,7-Tetramethoxythianthrene, 71b, also shows two reversible one-electron oxidations even under conditions in which the second oxidation of thianthrene, 71a, is irreversible owing to the greater stability of the dication of 71b [184,185]. However, more recent studies have provided a more nuanced picture of the redox chemistry of these species. Using fast scan voltammetry at reduced temperatures, reversible dimerization of the radical cations of 71a and 71b has been detected [186]. Furthermore, in  situ UV–vis and EPR/UV-vis-NIR spectroelectrochemical studies provide evidence for dimer formation that depends on temperature, concentration, solvent, and supporting

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Sulfur-, Selenium-, and Tellurium-Containing Compounds

electrolyte [187]. The dimer is diamagnetic as shown by EPR studies, but its structure in solution has proven controversial. In the solid state (71a)+ • AlCl4− [188] and (71b)+ • I3− [189] form dimers 72a and b, respectively, with planar heterocyclic rings. 2+ X

2+

X

S

S S

X X

S

X S

S

X

S X

S

X

72a, X = H b, X = OMe

73

However, on the basis of PM3 calculations, it was suggested that 71a+ • forms σ-dimer 73 with bent heterocyclic rings [186], but it has been argued the low −ΔG value for dimerization in solution favors π-dimer formation [190]. Near IR spectroscopic absorbance has also suggested the presence of a charge-transfer complex of thianthrene 71a with its corresponding radical cation [187]. Thianthrene 71a has also been reported to function as a redox catalyst in the oxidation of 2,3-dimethyl-2-butene [191], guanosine and DNA in acetonitrile [192]. Thianthrene has also been shown to participate in electrogenerated chemiluminescence [193]. Here the idea is that electrochemically generated thianthrene radical cation reacts by electron transfer with an electrochemically generated radical anion, such as that generated from 74 [194] or carbon dioxide radical anion [195] (produced by oneelectron oxidation of oxalate by thianthrene radical cation). S S+

S+ N

S

N O

S +

S CH2COp–MeC6H4

74

75

76

Owing to the energetics and rapidity of the electron transfer an excited state species is formed that luminesces. Chemiluminescence is also achieved by injection of hot electrons from oxide coated metal (Ta, Pt) electrodes into thianthrene radical cation [196,197]. The reactions of thianthrene radical cation with nucleophiles have generated considerable interest [88]. An overview of such chemical studies has recently been published [198]. Surprisingly the reaction is second order in radical cation, which suggests a disproportionation mechanism followed by rapid reaction of the dication with water. However, electrochemical studies [199–201] revealed that water first associates with the radical cation and this complex undergoes rate-determining electron transfer with another radical cation. The nucleophile-radical cation is a π-complex devoid of a covalent bond between nucleophile and radical cation until electron transfer occurs (complexation mechanism). In the related half-regeneration mechanism, a bond is formed between nucleophile and radical cation. Pulse-electrolysis stopped-flow methods have been used to determine the kinetics of the reactions of 71b radical cation and dication with methanol and pyridine [202,203]. As expected,

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Organic Electrochemistry

the dication reacts faster with nucleophiles than the radical cation. With methanol the reactions of  the dication and radical cation are both second order in methanol, but first order in dication and second order in radical cation. Thus, in both cases, complexes of the ions with methanol must be deprotonated by another methanol molecule in (dication) or preceding (radical cation) the ratedetermining step. With pyridine the dication reacts directly in the rate-determining step, resulting in first-order dependence in dication and pyridine. With the radical cation, the reaction is second order in radical cation and first order in pyridine. Hence, the rate-determining step involves the well-precedented electron transfer from the radical cation complex to another radical cation. The reactions of thianthrene radical cation with nucleophiles are not only of mechanistic interest but of synthetic interest as well. The electrosynthesis of 75 and 76 have been reported [204] by controlled potential electrolysis in acetonitrile of thianthrene with 4-methylacetophenone and cyclohexene, respectively. The electrochemistry of other thianthrene derivatives and analogues has been investigated. The electrochemistry of thianthrene appended with one or two ferrocene rings have been reported [205]. Cyclic voltammetric studies of 77a and b in dichloromethane with Bu4N+ [B(C6F5)4]− as supporting electrolyte shows two reversible oxidations for 77a with E1/2 = +0.33 and +1.32V and three reversible oxidations for 77b with E1/2 = +0.27, +0.44, +1.55 V versus Ag/Ag+. Y

X S

S

S

S 2

77a, X = Fc; Y = H b, X = Y = Fc

78

The first peak for 77a and the first two peaks for 77b are ascribed to one-electron oxidation of the ferrocene moieties and the peak at highest oxidation potential for both is ascribed to one-electron oxidation of the thianthrene moiety. The separation of the two ferrocene oxidations in 77b indicates interaction between these moieties. 1,1′-Bithianthrene 78 shows two not totally separable two electron oxidations in acetonitrile [206]. Cyclic voltammetric studies on thianthrenes substituted with heterocycles and bridged with heterocycles have been reported [207]. Cyclic voltammetric studies in acetonitrile or dichloromethane of thianthrene derivatives 79a and 79b and c show one or two reversible oxidation potentials, respectively [206,208], and 80 exhibits a quasi-reversible oxidation [208]. Et

Et

Et

Et R

S

R

R

S

S

R΄

S

R΄

R

S

S

Et

Et

Et

Et

79a, R = R΄= H b, R = R΄= SC8H17

Et

Et

c, R = R΄= SCH2CH2S

When coated on a glassy carbon electrode, 81 shows a reversible oxidation [209]. S S

S 81

© 2016 by Taylor & Francis Group, LLC

R

S Et

80, R = SC8H17

S

R

S

Et

1055

Sulfur-, Selenium-, and Tellurium-Containing Compounds

VI.

TETRATHIAFULVALENE AND ANALOgUES

Interest in electron transfer from TTF, 82, began with the report that it underwent two reversible one-electron oxidations at relatively modest potentials [210]. This interest was spurred on by the discovery of unexpectedly high conductivity in its charge transfer complexes [211], thereby earning the sobriquet of “organic metals” (although one-dimensional conductors) for these complexes and even superconductivity in so-called “Bechgaard salts” [212,213].

S

S

S

S

S

S

S

S 82

83

Interest in hybrid salts: TTF radical cation-inorganic anions with unpaired spin (in part to couple conductivity with magnetic properties) [214,215], single-component molecular metals [216], covalently linked substituted TTF and TTF analogues with electron-acceptor moieties [217,218] and with each other [219] have blossomed. An important driver in this research is the development not only of organic conductors and superconductors [220,221] but organic field effect transistors [222], nonlinear optical materials [223,224], molecular rectifiers [224], chemical sensors [224,225], redox switchable ligands [224–226] for molecular shuttles, switches [224], and photovoltaic and solar cells [227]. A brief perspective on “organic metals” and applications to other areas of molecular electronics has been presented [228]. To monitor drugs, appending them with an electroactive group is attractive and a recent example involves affixing TTF to benzodiazepines [229]. This update from previous reviews [1,230] will be selective rather than comprehensive and focus on electrochemical studies of TTF, its derivatives, analogues—especially its π-extended analogues 83 and the significance of these studies. Some recent studies on some substituent effects on the oxidation potentials of substituted TTF have been reported. The first oxidation potentials for 84a and b and 85a are comparable to those of 84c and 85b, respectively, but the second oxidation potentials are moved to more anodic potentials [231]. Y Y

X

S

X΄

S

S

S

S

S

X Y

Y΄ X

S

S

X΄

84a, X = Y΄ = S; X΄ = Y = CH2 b, X = X΄ = CH2; Y= Y΄= S c, X = X΄ = S; Y= Y΄ = CH2

Y΄

X

Y

85a, X = S , Y = CH2 b, X = CH2 , Y = S

Both first and second oxidation potentials are rendered more anodic for 86 than 84c [232]. Both oxidation potentials are rendered more anodic with electron-withdrawing substituents (in monoaryl) and less anodic with electron-donating substituents in the arene ring of monoaryl monocyclic TTF 87a relative to TTF, 82.

© 2016 by Taylor & Francis Group, LLC

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Organic Electrochemistry S

S

S

S

S

S

S

S

Ar

S

S

Ar΄

Ar΄

S

S

Ar΄

S S

S

86

87a, Ar΄= H b, Ar = Ar΄

The effect is even larger with tetraryl TTF, 87b [233]. Other recent electrochemical studies provide additional information on subsituent effects on TTF oxidation potentials: esters, phthalimides substituted benzo TTF [234], 1,4-dithiin [235], and thiophene, substituted dibenzo TTF [236]. Thiophene annulated TTF 88 has attracted particular interest because of its high mobility in field effect transistors [237,238]. Cyclic voltammetry of 88 in DMF shows two reversible 1e− oxidations with E1/2 = +0.78 and +0.96 V versus SCE [237]. C6H13

C6H13 S

S

S

S

S

S

S

Me

S

S

S

S

88

S

S

Me

S

S

2 89

Preliminary cyclic voltammetric studies on oligothiophene TTF derivative 89 have been reported, but interpretation of the complex results awaits further studies [239]. Additional substituent effects on the electrochemical oxidation of substituted derivatives of TTF have been reported [240–244] and an overview of such tuning for molecular electronics applications published [245]. A TTF derivative appended with two pyridine moieties was immobilized on a Pt electrode to determine the effect of solvent and air on its electrochemical oxidation [246]. A single pyrrolo-TTF molecule attached to gold in ionic liquid provided electrochemical potential control of its conductance [247]. Electrochemical responses to TTF moieties appended with metal binding sites have been reported [248]. Self-assembly of bis(pyrrolo)TTF appended with pyridine moieties on complexation with Pt to afford electroactive cages has been achieved and characterized [249]. Remarkably, the 4e− reduction of O2 to H2O is accomplished by TTF in acidic ClCH2CH2Cl and at an acidified water ClCH2CH2Cl interface [250]. Ion transfer voltammetry showed the transfer of 82 • + in this reaction. The proposed mechanism for the reaction is protonation of TTF, 82 to give 90 and then reduction of O2 to water generating 82 • + perhaps via TTF-90 aggregates. Electron transfer from TTF, 82, to 90 to generate 82• + has been proposed in hydrogen-bonding assisted self-doping of TTF salt 91 [251]. MeS S MeS

S

H

S S

S

S

MeS S

S

S

S

S

S

+

S

S 90

S

S 91

CO2–NH4+

MeS

S

S

SMe

92

The interaction of two TTF moieties especially on oxidation has attracted interest. Studies on TTF units linked by σ-bonds, conjugated π-systems, halogen atoms, or flexible chains have been reported in

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Sulfur-, Selenium-, and Tellurium-Containing Compounds

which through-bond or through-space interactions are possible [219]. Electrochemical and spectroelectrochemical studies provide evidence for such interaction. For example, the weak interaction between TTF• + and TTF, which is the basis for conducting Bechgaard salts in the solid state, is typically not observed in solution unless the moieties are constrained by covalent linkages. For example, the first two electron oxidation wave for 92, determined by cyclic voltammetry in benzonitrile, can be deconvoluted into two waves, suggesting interaction between one TTF• + unit with the other TTF unit (the effect is intramolecular because it is observed even on dilution) and NIR absorption at 1880 nm confirms the formation of the mixed valence species [252]. A tetracarbonyl-t-butyl calix[4]arene scaffold appended with two TTF moieties shows a broad first oxidation and a narrow second oxidation on cyclic voltammetric analysis. Chemical oxidation (NO+SbF6−) and spectroscopic studies show radical cation absorption and bands at 1750 and 765 nm ascribed to the mixed valence dimer (TTF)2• + and π-dimer (TTF)22+, respectively. Interestingly, addition of Na+ results in a new and sharp reversible redox system at more positive potential for the first wave, suggesting that Na+ complexation results in a conformation in which the two TTF moieties cannot interact [253]. Recently, interaction between two separate TTF molecules has been observed in a self-assembled coordination cage in which cyclic voltammetric behavior provides evidence for a (TTF)2+• mixed valence complex [254]. Two oxidation waves (200 and 500 mV vs. Ag/AgCl) are observed for oxidation to the radical cation and dication but square wave voltammetry shows the first oxidation wave to be split (152 and 304 mV vs. Ag/AgCl). Controlled potential electrolysis at a potential of 180 mV afforded the mixed valence complex (TTF)2+ •, which showed broad absorption in the near IR at 2000 nm. Further oxidation at a potential of 400 mV generated the radical cation π-dimer, (TTF)22+ [255,256], which escapes the cage due to electrostatic repulsion between cations as evidenced by the next oxidation peak potential that is comparable to that for the oxidation of free (TTF)• + to TTF2+. An analogous π-dimer identified by its visible absorption is formed in the cavity of cucurbit[8]uril [257]. Interestingly, cyclic voltammetric studies show that chiral TTF derivative 93, but not its meso-analog, shows that the typical first oxidation wave is split into two waves [258].

O

S

S

O

O O

S

S 93

This is interpreted in terms of association of the radical cation formed after 1e− oxidation with another molecule to form the mixed valent (TTF)2+ • in the chiral but not meso species. Near IR absorption at 2000 nm for this species supported this suggestion. Electrocrystallization of chiral dimethylethylenedithio TTF provided enantiopure and racemic cation dimers [259]. The enantiopure salts were semiconducting, but the racemic salt is an “organic metal.” Redox modulation of the chirooptical signal of chiral TTF fused to helicenes has been accomplished [260]. The observation of overlapped first oxidation waves in compounds 94, 95 suggests interaction between TTF units, but this needs to be clarified [261]. R΄ RH2C

BuS CH2R

R=

S BuO

BuS R΄

S

S

R΄ CH2R

RH2C

S S

SMe

2

RH2C

R΄ = Me, Et, OMe 94

95

There are other recent examples using electrochemistry that demonstrate the lack of intramolecular interaction between TTF units. Thus, two reversible waves are found for 96 by cyclic voltammetry

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1058

Organic Electrochemistry

in PhCN as with TTF, but the first oxidation potential for 96 occurs at a more positive potential than that for TTF (E11/ 2 = 0.23 V, E12/ 2 = 0.57 V and E11/ 2 = −0.07 V, E12/ 2 = 0.48 V vs. Ag/AgNO3 in PhCN, respectively) [262]. Thus, there is no interaction between TTF units in 96, but the electronwithdrawing acetylene substituents render the first oxidation more difficult. The smaller ΔE for 96 than for TTF (0.34 vs. 0.55 V) is cited as evidence for greater charge delocalization in 96 on oxidation than for TTF and consequent reduction of Coulomb repulsion. Cyclic voltammetry of 97–99 in CH2Cl2 shows two redox couples in the range of E11/ 2 = 0.23 − 0.51 V and E12/ 2 = 0.84 − 0.90 V, versus SCE [263]. This is comparable to a model mono-TTF derivative (BEDT-TTF), which shows two redox processes with E11/ 2 = 0.48 V and E12/ 2 = 0.89 V versus SCE under the same conditions, demonstrating that this is no intramolecular interaction between TTF units in 97–99. Bu R

S

CO2(CH2)2R

S

R

CH(CH2CH2R)2

R

Bu S

R=

R

S

CH(CH2CH2R)2

S

S

97

96

SR΄

S

S

S

R=

S

SR΄

S

R΄, R΄= ( CH2)2

R΄= CH3 98 R

R

CO2(CH2)2R

99

R SBu

NC(CH2)2S

S

R=

S

S

CH2S R

S

R R

R 100

CH2R CH2R

RCH2

SMe

MeS S

R = CH2R

RCH2

MeS

CH2R 101

© 2016 by Taylor & Francis Group, LLC

S

S S

S

SBu

1059

Sulfur-, Selenium-, and Tellurium-Containing Compounds

Similarly, a variety of TTF moieties appended to the triptycene, for example, 100 showed only two reversible oxidations by cyclic and differential pulse voltammetry in CH2Cl2 [264]. Thus, the TTF moieties oxidize independently, a conclusion that is also supported by spectroelectrochemical studies. The TTF units in 101 also do not interact electronically as evidenced by only two oxidation peaks (the first of which was shown to be a 6e− oxidation) in cyclic voltammetry studies [265]. Cyclic voltammetric studies in PhCN reveal no interaction between the pyrrole fused TTF moieties in 102–104 because two reversible oxidations are observed [266]. Although analogue 103a shows a broad first oxidation wave and analog 104a shows a split oxidation wave, these results are not ascribed to intramolecular interaction because they are not found for 102a, but rather to intermolecular interaction. R R

R

R

R

R

R

R R

R

R

R

R

SR΄ R=

N

S S

102

S SR΄

S

103

a, R΄= nBu b, R΄= nC12H25

104

In addition, the redox chemistry of a bis-TTF calix[2]pyrrole[2]thiophene shows that the two TTF moieties do not interact. Cyclic voltammetry in CH2Cl2 shows two reversible 2e− oxidations at the same potential as that for the mono TTF model, indicating that the 1e− oxidation of both TTF moieties occurs at the same potentials [267]. Similarly, cyclic voltammetric studies of a bis-TTF calix[4]arene in which the TTF moieties bridge the “upper rim” also demonstrate that there is no electronic communication between the TTF moieties [268]. In addition to calixarenes bearing two TTF moieties, TTF bridging two calix[4]arenes has been reported and studied electrochemically [269]. Cyclic voltammetric studies in CH2Cl2 –CH3CN (1:1, v/v) show the expected two redox processes with the first reversible (E11/ 2 = 0.43 V vs. Ag/AgCl) but the second at Epox = 1.07 V is broadened and quasi-reversible for unknown reasons. Interestingly, TTF dendrimers have been studied electrochemically and the TTF units act independently and are all oxidized, as ascertained quantitatively by current measurements, at the same potentials, resulting in multiradical cations in the first oxidation step and multidications in the second step [270]. In the presence of nonredox active electron-rich oligomers, complexation with oxidized TTF oligomer has been inferred from the substantial decreases in the redox currents compared with that in the absence of the electron-rich complexing oligomer owing to a decrease in the diffusion rate of the complexed species [271]. A review on these mixed valence complexes and related systems has recently appeared [272], and the significance of splittings in electrochemical oxidation of such systems discussed. Interaction between TTF and triarylmethyl radicals linked by an ethylene moiety has been evaluated electrochemically [273]. Dimerization between two different TTF moieties can be controlled in redox active [2]catenanes [274]. π-Extended TTF analogues, exTTF, in which two 1,3-dithiolenes are separated by a conjugated π-system have attracted much interest and show unusual electrochemical behavior [275]. Cyclic voltammetric studies of TTF vinylogues 105 show that substituents affect the oxidation potentials but, in addition, increasing the conjugation length (i.e., increasing n) lowers the first oxidation

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1060

Organic Electrochemistry

potential and decreases the difference between the first and second oxidation potentials [276]. Even more interesting is the behavior of 106 that is synthesized by controlled potential electrolysis of 1,4-dithiafulvalene 107 [277]. R Me

S R

S S

S

S

Me

SMe

S MeS

n

Ar

R

S

S

R

S

Ar

105

MeS

Me

106

S

Ar

107

Electrochemical studies show that the mechanism for this synthesis involves the dimerization of the radical cation of 107. Two reversible 1e− oxidations or one 2e− oxidation occurs depending on the Ar substituents, their positions (o– vs. p–), and the solvent [277]. These results are rationalized in terms of electronic and steric factors in which there is a large conformational change on oxidation [277,278] underlying the potential inversion [279] and consequent a single 2e− oxidation wave. This analysis is supported by X-ray structural studies and DFT computations. Cyclic voltammetric studies on 108 provided insight into its unusual redox behavior [280]. In CH2Cl2, 108 undergoes two 1e− oxidations followed by a 2e− oxidation to afford the corresponding tetracation 1084+. However, on the reverse scan at scan rates greater than 100 mV/s, in addition to reduction waves corresponding to the oxidations, an additional more negative reduction peak is observed. On the basis of theoretical calculations, the nonplanar s-trans isomer (about bond a,b) is favored for 108, but the nonplanar s-cis isomer is favored for 1084+. Consequently, at fast scan rates the reduction of s-cis-108+• can be observed in addition to that of s-trans 108+•. The effect of aromatic rings connecting 1,3-dithiafulvalene units has been reviewed [275]. More recent studies have probed the consequences of such connectors on the use of such compounds as organic field effect transistors [281]. The redox chemistry of one such system connected by a benzene ring but with thiophene appended as well 109 has been reported [282]. SMe

MeS S Me Me S Me

b a

C8H17

S

S S

S

Me

S

C8H17 S

S 108

S

S SMe

MeS 109

Notably the oxidation is dominated by the exTTF core and electropolymerization of the thiophene moiety does not occur. Electrochemical studies on 110, in which thiophene bridges 1,3-dithiafulvalene and exTTF moieties, have been reported in PhCN as solvent [283]. Deconvoluted cyclic voltammetry shows four pairs of reversible redox waves. The first two waves are 2e− processes and the following two 1e− processes. p-Quinodimethane TTF analogue 111 undergoes two reversible 1e− oxidations at −0.11 and −0.04 V versus SCE in acetonitrile, which is much less anodic than that for the oxidation of TTF (0.28 and 0.64 V vs. SCE) under the same conditions [284] owing to reduced Coulombic repulsion on oxidation. However, analogue 112a undergoes 2e− oxidation at a potential slightly more positive

© 2016 by Taylor & Francis Group, LLC

1061

Sulfur-, Selenium-, and Tellurium-Containing Compounds

than the first 1e− oxidation for TTF (0.44 and 0.37 V, respectively, in CH2Cl2 vs. SCE) [285]. Thus, there is potential inversion; that is, the radical cation of 112a oxidizes at a lower potential than 112a. Cyclic voltammetric analysis in DMF led to an estimated potential inversion for 112a of 0.28 V [286]. Similar results for benzo-annulated analogues of 112a have been reported [285]. Thus, the radical cation of 112a appears to be destabilized relative to the dication and indeed, the radical cation of 112a generated by pulse radiolysis [287] or the radical cation of 112b generated by flash photolysis [288] disproportionates. R

R R

R

S

S

S S

S

R

R

R

R

S

S

S

S

S

S

S

S

S

S

R

R

S S

S

R

R

S

S

R

R

110a, R = SMe b, R = Me

111

112a, R = H b, R = SMe

The role of aromaticity in potential inversion is illustrated on electrochemical studies comparing 113 with 114. S+

S

+

S

S

S

S

113

S

114

S 115

Potential inversion is found for 114 but two 1e− oxidations are found for 113 [289]. Potential inversion is ascribed to aromaticity of the dication 115 obtained from 113. The importance of substantial geometry change on oxidation of 112, in which the saddle-shaped central ring of 112 becomes planar on 2e− oxidation with almost perpendicular 1,3-dithiolium substituents has been well illustrated by molecular constraints that prevent planarization [290]. S

S S

MeS

S

(CH2)5

S

SMe S

(CH2)5

S

S

S

S 116

S MeS

S

S

S

S

SMe SMe

S SMe 117

Thus, doubly bridged derivative 116 shows two reversible 1e− oxidations in CH2Cl2 at more anodic potentials than the 2e− oxidation for comparable nonbridged analogues [290].

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1062

Organic Electrochemistry

Electrochemical studies were used to evaluate electronic interactions between a TTF and an exTTF moiety linked by π-systems. Cyclic voltammetry of 117 in CH2Cl2 revealed three reversible peaks with oxidation potentials of 0.27, 0.71, and 1.12 V versus SCE [291]. The first peak corresponds to a 2e− oxidation and was assigned to the exTTF moiety. Cyclic and square wave voltammetry of 118a–c were solvent dependent (THF, CH2Cl2, and CH3CN were used as solvents) and showed a 3e− oxidation, for which separation of oxidation peaks was only observed for 118a in THF, followed by a 1e− oxidation [292]. Thus, the 2e− oxidation of the exTTF moiety and the 1e− oxidation of the TTF moiety overlap (except for 118a in THF) followed by further 1e− oxidation of the TTF moiety. R

R

S

S S S

S

S

R

R

S

S

S

S

S S

S

S

S

SMe

SMe

S

S 118a, R = H b, R = SMe c, R, R = (SCH2)2

S

S

S

119

Comparison with models with either exTTF or TTF moieties but not both showed features ascribable to each moiety in 118 and also electronic interations between them. However, far more consequential electronic interactions between exTTF and TTF moieties were demonstrated by electrochemical studies of 119 in which the moieties are directly connected without an intervening π-system (unlike 117). Here, deconvoluted cyclic voltammograms with CH2Cl2 as solvent showed four separate reversible oxidations [293]. Compounds with exTTF moieties attached by unstrained saturated linkages as illustrated by 120 do not interact electronically. Consequently, cyclic voltammetry shows a typical quasireversible oxidation peak for all of the exTTF moieties, which for 120 results in a 6e− oxidation [294]. SMe

SR Me

S

S

S

S

Me R=

RS

SR Me

MeS

SMe

120

Likewise cyclic voltammetry studies show one pair of quasi-reversible waves for the noninteracting 1,3-dithiafulvalene moieties in 121 [295]. However, for compound 122 in which two exTTF moieties are directly linked by a σ-bond, two closely spaced 2e− oxidations are observed electrochemically ascribed to separate 2e− oxidations of each exTTF moiety [296]. This interpretation is supported by spectroelectrochemical studies and DFT calculations.

© 2016 by Taylor & Francis Group, LLC

1063

Sulfur-, Selenium-, and Tellurium-Containing Compounds

RS

S

S

S

S

RS

SR SR

Me

S

S

Me

S

S

Me

2 121, R = Me, Et, Bu

122

As pointed out in the introduction to TTF and analogues, interest in these systems is driven by their applications. The following text is organized in terms of these applications, but the emphasis is on the electrochemistry of the organosulfur components. A review on TTF and exTTF vinylogues as building blocks for organic materials and redox chemistry control of their conformational switching has been published [297], and a review of the application of exTTF moieties in organic electronics has appeared [298]. Compound 123 has been shown to be a selective chemiluminescent probe for singlet oxygen. MeS S MeS

S(CH2)2O

S

S

S

S

S

S

S

O CONH(CH2)2NHCAr

124, Ar = 3,4,5–(C12H13O)3C6H2

123

Cyclic voltammetric studies in CH3CN show three oxidation waves at 0.47, 0.80, and 1.3 V versus SCE [299]. Stepwise electrochemical and chemical oxidation of TTF serves as input signals and the consequent distinct absorptions for the cation radical and dication as output signals in a molecular logic gate (half-adder) that, owing to the facile reduction of TTF oxidation, is resettable [300]. Compound 124 shows oxidation potentials shifted to less positive values when assembled as a gel than in PhCN solution ascribed to π-stacking in the assembled gel [301]. Compound 125a shows two 1e− oxidations by cyclic voltammetry and oxidized xerogel from 125a is a semiconductor [302]. On the other hand, electrochemical oxidation of the gel formed from 125b in ClCH2CH2Cl destroys the gel by impairing intermolecular hydrogen bonding by the urea moiety [303]. Oxidation of end functionalized TTF poly (N-isopropylacrylamide) micelles also disrupts the micelles [304]. S S S

S

S

SR

S

R΄

125a, R = CH2CONHC18H37 R΄= SMe b, R = (CH2)2NHCONHC12H25 R΄= H

S S

S

Ar

S

126a, Ar = pNO2C6H4 b, Ar = pNCC6H4 c, Ar = 4-pyridyl d, Ar = 4(NMe)pyridinium

Chiral TTF compounds have attracted interest because of the possibility of synergism between conductivity due to the TTF moiety and chirality due to the chiral moiety [305,306]. Chiroptical molecular switches have been reported [307,308]. For example, cyclic voltammetry of a chiral polymer containing TTF moieties shows two quasireversible waves that enable the production of different oxidation states that display different CD effects [308]. Conjugation of TTF moieties with π-electron-accepting moieties has attracted interest for their nonlinear optical properties. Here, intramolecular charge transfer is evaluated electrochemically. Thus, cyclic voltammetry shows anodic shifts for the two 1e− oxidations in 126a and b [223], 126c and d [309], 127 [310], 128a–c [311], and analogues. π-Extended donors with TTF moieties attached to a vinyldithiolene show two reversible oxidation waves [312].

© 2016 by Taylor & Francis Group, LLC

1064

Organic Electrochemistry

Molecular motions controlled by the redox state of TTF moieties forming the basis of molecular tweezers, clips, and switches have garnered much attention recently. R S

S

S

S

CN n

S

R

CN

Ar

S

S

S

Ar

R = MeS, H, S(CH2)2S 128a, Ar = 2–pyridyl

127

b, Ar = 3–pyridyl c, Ar = 2–quinolyl

Cyclic voltammetric studies on 129 in CH2Cl2 show that there are two 1e− oxidation waves followed by a reversible 2e− oxidation [313]. OR RS S RS

S

S

S

S(CH2)nS

S

S

S S

S

S

S

SR

S

SR

OR

S

S

S

S

Ar

H

H

Ar

129, R = C6H13

130

Thus, the conformation of 129 changes from closed, where the radical cation of one TTF unit interacts with the other TTF unit, to open owing to electostatic repulsion between TTF dication units [314]. Anodic oxidation of 130 resulted in intramolecular coupling of the 1,3-dithiafulvalene moieties via the cis-radical cation to give 131 dication that is reduced to 131 [315]. SMe

MeS S

SMe

MeS

S

S

SMe SMe

S

S

S(CH2)nS S

S

R= S

S S MeS Ar

Ar

R

R

131

132 S

S

S

S

+

+

R

R

R=

133

© 2016 by Taylor & Francis Group, LLC

S

S

S

S

N+

+

S SMe

OCH2CH2O

134

4

S

Sulfur-, Selenium-, and Tellurium-Containing Compounds

1065

Whether two 1e− or one 2e− process occurs depends on the chain length, n, and the Ar substituents. With a short link, compound 131 acts as reversible molecular clips that close on oxidation [315]. Other cis-locked analogues have been studied recently [316] as has compound 132, which acts as an electrochemically activated tweezer [317]. Cyclic voltammetry of 132 in CH2Cl2 shows only one oxidation peak at +0.68V corresponding to a 6e− oxidation (resulting from oxidation of the bis-tetrathiafulvalene and both exTTF units) but on scanning cathodically two reduction peaks are observed at +0.55 and +0.23 V versus Ag/AgCl [317]. Electrochemical data have been cited as evidence for conformational change in TTF-bridged resorcin[4]arene cavitands [318]. In this molecule with two TTF moieties, cyclic and differential pulse voltammetry show a broadened 2e− reversible step followed by a sharper second 2e− reversible oxidation. It is argued that in the first oxidation step the conformation is such that the two TTF moieties are close enough to interact but the oxidized TTF moieties repel each other electrostatically, resulting in a conformational change in which the two moieties do not interact with each other and further oxidize in the second step. Pseudorotaxanes, rotaxanes, and catenanes containing TTF units have been studied as redox activated switches [319] (see also Chapter 12). Electrochemistry is used to probe the structures of these materials and their precursors and, in particular, electronic interactions affecting TTF redox chemistry. For example, an acyclic polyether with a central TTF moiety threads through cyclophane 133, resulting in attractive interactions between the electron-rich TTF and electrondeficient bipyridinium moieties. This results in an anodic shift for the first oxidation but not the second oxidation and scan rate dependence of the first but not the second oxidation of the TTF moiety [320,321]. The first oxidation is rendered more positive by the donor–acceptor interaction, but the second oxidation occurs after dethreading of the TTF polyether. Interestingly, the dethreading is reversed on reduction, resulting in electrochemical control of the threading/dethreading process. In another example, a rotaxane is formed in which a TTF polyether is threaded through an α-cyclodextrin cavity and stoppered (larger groups are appended to the ends of the TTF polyether) [322]. Cyclic voltammetry shows two reversible 1e− oxidations for the TTF moiety in this rotaxane with a shift in oxidation potential for the first but not second process (0.32, 0.55 and 0.17, 0.54 V, respectively, vs. SCE). Thus, the electrochemical evidence suggests that the α-cyclodextrin encloses the TTF moiety but on oxidation of the cyclodextrin moves away from the TTF radical cation. This interpretation is supported by spectroscopic results. Reduction reverses this process. Electrochemical studies of an analogous system with cholesterol stoppers that enable gel formation have been reported [323] as a well as another rotaxane capable of forming liquid crystals in which rearrangement occurs on oxidation and is reversed on reduction. These redox controlled movements change the liquid crystal ordering resulting in electrochromism [324]. [2]Catenanes are formed by interlocking 133 with a cyclic polyether with a TTF moiety and a 1,4-dioxybenzene or 1,5-dioxynaphthalene moiety transannular to each other. Here, the first oxidation of the TTF moiety is anodically shifted and scan rate dependent and results in rearrangement owing to the Columbic repulsion [321]. The rearrangement involves rotation resulting in removal of the TTF radical cation moiety from the cyclophane cavity and its replacement by the dioxyaromatic unit. A [3]catenane in which two cyclic polyethers each containing a TTF moiety are interlocked with cyclophane 133 gave interesting electrochemical results. Four separate 1e− oxidations occurred, indicating interaction between TTF units and, of particular interest, mixed valence (TTF)2+• and radical cation dimer (TTF)22+ formation within the 4+ cyclophane despite electrostatic repulsion  [325]. A  pseudorotaxane consisting of 134 threaded by dibenzylammonium salts has been reported [326]. Cyclic and differential pulse voltammetry on 134 shows two closely spaced oxidations, indicating that oxidation of one exTTF unit influences the oxidation of the other. Other examples of switches and other TTF and exTTF-containing molecules designed for switches and molecular electronics and their electrochemistry have been reviewed [327]. Use of TTF-based molecules as electrochemical sensors for metal cations has been reported. Thus, 135 shows two reversible oxidations. On addition of Pb2+, the first oxidation shifts anodically but not the second.

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1066

Organic Electrochemistry MeS S

S

RN S

S

NR RS

S

OCH2

CH2OCH2

CH2O(CH2)2

OCH2

CH2OCH2

CH2O(CH)2

S

R= 2

S

S S

S

O

O

OMe

O

O

OMe

136a, R = Me b, R =

O(CH2)6

O

135

O

S

n

The shift in the first oxidation is due to complexation by the crown moiety in 135. The peak potential for the Pb2+ complex of 135 is 140 mV more anodic than that for 135. On oxidation of the lead complex, the Pb2+ is released, resulting in no shift in potential for the second oxidation. In addition, Ba2+ complexes also result in anodic shift of the first oxidation of 135 by 90 mV [328]. Addition of Pb2+ to 136a gave analogous results, that is, owing to metal complexation anodic shift of the first but not the second oxidation potential because the metal ions have been released. Furthermore, polymeric films of 136b show similar behavior as 136a enabling redox control of Pb2+ binding to the surface [329]. However, appending TTF vinylogues with polyether metal cation complexation sites results in the unusual feature that Ba(ClO4)2 and Pb(ClO4)2 complexes more strongly associate with the dication than monocation despite the increase in charge [330]. This is ascribed to a conformational change on oxidation (clip motion) and conformational switching with TTF vinylogues with naphthyl substituents [331] has also been reported. Dendrimers featuring 137 units electrodeposited on the electrode surface also show redox dependent complexation of Ba2+ and, in particular, anodic shift of the first oxidation in potential on addition of Ba2+ [332]. Electrochemistry shows a modest but analogous effect with a bis (calix crown) TTF species and Na+ in solution [333]. RS S S O

SMe

S

S

S

S O

O

O

O

S

O

O

S

S

S

S

S

S

S

O

O

O

O

3 137

138

Cyclic and square-wave voltammetric studies show a quasi-reversible 2e− oxidation of 138. Addition of metal cations results in anodic shift of the first 2e− oxidation due to metal complexation with the crown ether moieties. The largest measured shift of 115 mV was found for Ag+ and the shifts follow the order Li+ < K+ < Ba2+ < Na+ < Ag+ [334]. Compounds with TTF moieties appended with metal binding sites more remote from the redox moiety but conjugated with it have been reported. Cyclic voltammetry of 139a and b in CH3CN shows two reversible waves at more anodic potentials than TTF owing to charge transfer from the TTF moiety to pyridine moiety. X S S

S

N

S S

S 139a, X = b, X = c, X =

© 2016 by Taylor & Francis Group, LLC

X

CH CH C C C N

S S

140a, X = CH b, X = N

N

1067

Sulfur-, Selenium-, and Tellurium-Containing Compounds

This charge transfer is increased on addition of Pb2+, which coordinates to the pyridine moiety, resulting in spectroscopic and electrochemical changes. In particular, both oxidation potentials are rendered more anodic, demonstrating that, in contrast to the preceding examples, the metal ion is not liberated on TTF oxidation [335,336]. Similar spectroscopic and electrochemical results were reported for Pb2+ addition to 139c [337] and electrochemical oxidation potential shifts for 140a on addition of Cu2+ [338]. Compound 140b shows two 1e− oxidation waves anodically shifted from those of TTF owing to the electron-withdrawing pyrazine moiety. Further anodic shifts of both oxidation waves are seen on metal coordination, for example, Mn(hexafluoroacetylacetonate)2. In this case, cyclic voltammetric studies over time enable the monitoring of dissociation of the Mn(II) complex [339]. Redox potentials for a TTF derivative fused to a pyrazine ring have been reported [340]. Compounds with two TTF moieties bridged by a pyridine moiety, for metal coordination, have been prepared. In this case, coordination of a magnetic metal to obtain conducting magnetic materials rather than metal ion detection is the goal, but the ligand electrochemistry is of interest. Thus, 141 shows two 1e− oxidations, separated by 166 mV, and a 2e− oxidation.

MeS MeS

S

S

S

N

S

S S

S S

S

S

N

SMe

S

N H

S

HN

SMe

S

S

SC5H11

NH S

H N

141

S

SC5H11

142

The observation of two 1e− oxidations (each TTF unit oxidizing at a different potential) rather than a 2e− oxidation suggests interaction between the two TTF moieties forming π-dimers (owing to ring flexibility) on oxidation [341]. In addition to electrochemical sensors for metal cations, anion sensors based on TTF have been reported as well. Thus, addition of Cl− or Br− to 142 in CH3CN results in a cathodic shift in its first oxidation potential, reaching a limit of about 40 mV with a stoichiometric amount (1:1) of halide [342]. In a related example, an analogue of 142 with four benzo-TTF moieties fused to the calix[4]pyrrole core forms a molecular complex with Li+@C60. On addition of Cl− electron transfer from TTF to C60, moieties occurs based on cyclic voltammetric analysis in which the 1e− oxidation potential of the TTF moiety shifts anodically by 130 mV [343]. Cyclic voltammetric studies of 143 shows that on addition of F− the first oxidation wave is shifted cathodically up to a maximum of 140 mV when 1.0 equiv of F− is added. C6H11S C6H11S

S

S MeS

S

S

N

S

CONH N

2 143

S

B MeS

S

S 144

The shift is ascribed to binding of F− to B that reduces the electron withdrawal by B [344]. Another sensor with TTF attached to a boron-dipyrromethene also shows a 17 mV cathodic shift of the first oxidation potential of the TTF moiety [345]. Larger electrochemical shifts in reponse to F− have been reported for another TTF-based sensor [346]. Addition of AcO −, F−, Cl−, or Br− to 144 resulted in no changes in its cyclic voltammogram, but addition of H2PO4− resulted in two new oxidation waves cathodically shifted from those in its absence by 100 and 158 mV. Complexation of 144 with H2PO4− renders it easier to oxidize and the complex has a 2:1 stoichiometry (144 to H2PO4−). Thus, 144 serves as a selective electrochemical

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1068

Organic Electrochemistry

sensor for H2PO4− [347]. Compound 145 also shows a cathodic shift in its first oxidation potential of its TTF moiety on addition of H2PO4− of up to 172 mV [348]. SO2NH(CH2)2NHCOR

C5H11

C5H11

S

S

NC

CN

NC

CN

SO2NH(CH2)2NHCOR S S

S

O Me

R= S

S

SMe

S

SMe

O

145

146

Cathodic shift of 112 mV of the first oxidation potential is found by addition of 2 equiv. of H2PO4− to a sensor with TTF fused to a diindolylquinoxaline [349]. Curiously the second oxidation potential in this system undergoes an anodic shift on addition of H2PO4−. Analogous behavior has been noted in other systems [345], but the most dramatic manifestation of anodic shifts in TTF oxidation on addition of anions is manifested in a calix[4]arene appended with four TTF moieties [350]. This compound shows unusual electrochemical behavior in that four redox processes are found by cyclic voltammetry, suggesting interaction of the TTF moieties on oxidation. Furthermore, addition of anions results in anodic not cathodic shift of all of the oxidation peaks. Addition of H2P2O72− results in a dramatic shift of the fourth oxidation of up to 453 mV. Clearly binding of anions to TTF receptors may result in changes other than increasing the electron richness of the TTF moiety. Other redox sensors based on TTF have also been recently reported [351–353]. As referenced in the introduction of this section, the charge transfer complex formed from electron-donating TTF and electron-accepting TCNQ is highly conducting. Consequently, much effort has been expended on investigating systems in which TTF or analogues (D, donor) are covalently linked to an acceptor (A) in so-called D-A dyads. Furthermore, such compounds might serve as molecular rectifiers [354]. The DA link may involve σ-bonds, π-systems, or direct fusion. Electrochemistry has been used to evaluate D,A interaction in such systems and provide insight into the HOMO–LUMO gap as illustrated in the following selected examples. Electrochemical studies on 146 in which there is a flexible σ-linkage between TTF and TCNQ moieties show modest shift (20 mV) of the first oxidation potential of the TTF moiety and first reduction potential of the TCNQ moiety compared with the respective parent molecules [355]. This indicates little interaction between redox centers. However, their difference in potentials of only 170 mV equates to an exceptionally low HOMO–LUMO gap (which is substantially lower than that determined spectroscopically). The more rigidly σ-linked 147 exhibits two reversible 1e− oxidation waves that are shifted anodically by 200 and 60 mV, respectively, from the model TTF derivatives [356]. NC

CN

S NC S

S

S

S

S

CN N

S

SPr

S SPr

NC

N R

CN 2 147

NC

CN 148

In compound 148, the first and second oxidation potentials are anodically shifted from 148, R = H with more electron-withdrawing sulfonamide substituents 148, R = p-MeC6H4SO2, p-O2NC6H4SO2 [357].

© 2016 by Taylor & Francis Group, LLC

1069

Sulfur-, Selenium-, and Tellurium-Containing Compounds CN NC S

SMe

N

S

S S

S

S

SMe

O S

S

R΄

S

S

O

S S

SMe

150a, R = R΄= H b, R = H, R΄= Cl c, R = R΄= Cl

NC CN

SMe

R

149

In compound 149, a TTF moiety is rigidly connected via a π-linkage to a TCNQ-like moiety. Remarkably, electrochemical studies show little shift in the reduction potentials of the A portion but anodically shifted oxidation potentials for the TTF moiety [358]. The cyclic voltammetric data correspond to a HOMO–LUMO gap of 0.52 eV. DFT calculations show that the LUMO is localized on the TCNQ-like moiety and, although the HOMO interacts, it is substantially localized on the TTF moiety despite the fused π-connection between D and A. p-Benzoquinones have been used as the acceptor in D-A dyads with TTF. The first oxidation potentials (but, surprisingly, not the second) for the TTF moiety in the series 150a–c are shifted anodically as well as the first reduction potentials [359]. These shifts were ascribed to increasing charge transfer from TTF to benzoquinone moiety on going from 150a to b to c. Correlation of these redox changes with the spectroscopically measured charge-transfer band supports this suggestion (indeed both measure the HOMO–LUMO gap). Electrochemical studies on TTF linked to anthraquinone D-A dyads have also been reported [360]. Cyclic voltammetric studies have also been reported for A-D-A triad with a p-benzoquinone moiety fused to each end of TTF [361,362]. Here, the most interesting feature is that deconvolution of the cyclic voltammogram shows that the first reduction wave is split, indicating electronic communication in the mixed valence 1e− reduction product. Thus, the bridging TTF promotes electron transfer from the semiquinone to quinone moieties. Cyclic voltammetry was used to demonstrate no interaction between TTF and quinone moieties attached by a flexible polyether σ-linker until coordination by Pb2+, Sc3+, and Zn2+ ions that promote intramolecular electron transfer [363–365]. Similar electrochemical studies on TTF-naphthalenediimide dyads and DAD-triad σ-linked with a flexible polyether chain have been reported [366] as well as TTFperylenediimide-TTF triads [367]. Electrochemical studies with the rigid TTF-naphthalenediimide dyad 151 show three reversible oxidations, the first two due to the TTF moiety anodically shifted from a TTF model, and the third due to the 1,4-dihydropyrazine moiety [368]. C8H17 O

N

O H

Br Br O

N

N

S

N

S

S S

C5H11

S

C5H11

S

S

SPr

NO2

S NC

O H

C8H17

151

NO2 NO 2

O XC(CH2)2SO2

SPr

CN

152a, X = CH2O b, X = CH2N(Ph)

Linear and cyclic linked TTF-naphthalenediimides have also been studied electrochemically [366]. The thermodynamic driving forces for charge combination and charge-separation processes for a TTF moiety rigidly fused to a perylenediimide were calculated from electrochemically determined potentials [369]. Dyads 152a and b show two reversible 1e− oxidation waves due to the TTF moiety and three reversible 1e− reduction waves due to the A moiety. The first oxidation potentials are shifted anodically (40–70 mV) owing to electron withdrawal by the A moiety and both show small HOMO–LUMO gaps of ca. 0.3 eV [370–372].

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1070

Organic Electrochemistry

Anodic shift (0.10–0.15 V) in TTF moieties singly bonded to semisquarates relative to models provide evidence for charge delocalization [373] and the lack of such shift shows the lack of charge transfer in TTF moieties attached through several σ-bonds to perylene [374]. In compounds 153a and b, there are slight anodic shifts in TTF oxidation potentials compared with TTF and the somewhat broadened oxidation waves in 153b is suggestive of weak electronic interaction between the TTF moieties [375]. R R OMe N

N

N

R

N

S

S

N

R

N

N

R' N N

, R΄= OMe

153a, R = S

S

R

S

S b,

R

R

R = R΄= S

S

SPr

S

S 154, R,R = S

S

SPr

Connecting a TTF moiety via a flexible polyether chain to a flavin shows little interaction as evidenced by comparable oxidation potentials with a model, but on addition of Pb2+ or Sc3+ there is a dramatic shift in reduction potential of the flavin moiety as a result of metal coordination [376]. Annulating π-systems to connect a TTF moiety with a dipyridophenazine [377] or quinoxaline [378,379] moiety results in anodic shifts of their oxidation potentials. Anodic shifts in the oxidation potentials of the TTF moiety in an analogous compound with two quinoxaline moieties fused in an A-D-A fashion has been reported [380]. Thin layer cyclic voltammetric studies of 154 that features three TTF moieties π-linked through a hexazatriphenylene reveals that the first oxidation wave consists of three overlapping 1e− oxidations [381]. That is, each TTF moiety undergoes 1e− oxidation successively, but the second wave is very narrow, indicating that all three TTF radical cations are oxidized at the same potential to the corresponding hexacation. The TTF moiety in compounds 155a–c shows two reversible 1e− oxidations in CH2Cl2. On addition of HCl, both oxidation potentials are rendered more positive, but on addition of excess HCl the first oxidation is even more positive but the second oxidation potential is now about the same as that for the unprotonated species [382]. Linking an electron-rich methoxy thiazole by a σ-bond to TTF results in a cathodic shift in the two reversible 1e− oxidations of the TTF and the difference in these two oxidation potentials is solvent dependent [383]. Ph

R΄

R R

R΄ HN

N

Ph

H N

S

N

S

S

Ar S

155a, Ar = 2-pyridinyl b, Ar = 2-quinolinyl c, Ar = 6-MeO-2-quinolinyl

SPr

N

NH

SPr

R΄

R R

Ph S

156a, R =

R΄ SC5H11

S

S

S

S

S

SC5H11

R΄= H b, R = R΄=

S

© 2016 by Taylor & Francis Group, LLC

S

SC5H11 SC5H11

1071

Sulfur-, Selenium-, and Tellurium-Containing Compounds

Electrochemical studies of 156a in THF show two 1e− oxidations at −0.150 and −0.020 V versus Fc/Fc+, clearly indicating interaction between the two TTF rings mediated by the porphyrin ring [384,385]. Comparable interaction between the four TTF moieties in 156b is evidenced by the observation of at least five oxidation processes ascribable to the TTF moieties [384]. TTF moieties fused to porphyrins by quinoxaline linkers have been studied electrochemically [386]. Interestingly, a TTF moiety annulated to a porphyrin ring quenches porphyrin fluorescence by electron transfer from the TTF to porphyrin excited state but on oxidation of the TTF moiety to its radical cation, porphyrin fluorescence is no longer quenched [387]. Multielectron oxidations have been reported for TTF annelated to an expanded porphyrin [388]. Annulation of four TTF moieties to a phthalocyanine, like the porphyrin example, also results in interaction mediated by the phthalocyanine, but here the first reversible overall 4e− oxidation is resolved into two broad overlapping waves but not the second 4e− oxidation [389–391]. A related system was reported, but cyclic voltammetric studies were rendered difficult owing to aggregation in solution although two reversible oxidation waves were observed [392]. In contrast to the effect of TTF oxidation on the fluorescence quenching of a porphyrin outlined earlier, oxidation of TTF to its radical cation had no effect in a silyl phthalocyanine with a flexible link from the silicon to TTF moiety. That is, fluorescence quenching occurs whether or not the TTF is oxidized or not. It is suggested that quenching occurs in the oxidized TTF system by electron transfer from the TTF radical cation to the excited state phthalocyanine [393]. Neutral radical 157 in PhCN shows three reversible waves at potentials close to that of the compounds with constituent moieties. Use of PhCN/CF3CH2OH as solvent results in an anodic shift in the reduction potential of the oxophenalenoxyl moiety but no shift in the oxidation potentials of the TTF in 157 [394] constituent moieties. Electrochemical studies have also been reported for D-A dyads with TTF fused to pyridazine [395], benzothiazole [395], and directly attached to 1,3,5-triazine [396,397]. S

C8H17

S

S X

S

N

S S C60

O

S

O

S OC6H11

157

158

X=

S

S C6H11O

There has been much interest in D-A dyads with TTF or exTTF donors and C60 fullerene acceptors for use in solar cells. Those systems with exTTF donors have particular appeal because the lifetimes of the charge separated states exTTF+•/C60 −• can be unusually long (typically longer than TTF-C60 D-A systems) and consequently, may be especially serviceable in solar photovoltaic cells [227]. Again, here as before, electrochemistry is used to evaluate electronic interaction, or lack thereof, in the ground state, effect of substituents and determination of ΔGo for electron transfer. In an effort to control charge recombination in charge separated states, a dithienyl photoswitch was reported [398], applied to TTF-fullerene dyads [399] and both the open and closed form studied electrochemically. Electrochemical studies with TTF linked (saturated or unsaturated chain links) with pyrrolidine[60]fullerene show no significant interactions between moieties [400], as well as a comparable A-D-A triad and the redox data used to calculate ΔG o for electron transfer [401]. Similarly, electrochemical evidence reveals negligible electronic communication between exTTF and [60]fullerene moieties in compounds in which exTTF is covalently linked to pyrrolidine [60]fullerene [402] as well as comparable A-D-A triads [403].

© 2016 by Taylor & Francis Group, LLC

1072

Organic Electrochemistry

Similar analysis shows no significant interaction between exTTF linked through chiral binaphthyls to pyrrolidine[60]fullerene [404] as well as in a pseudorotaxane supramolecular structure with an exTTF derivative appended with a dibenzyl ammonium functionality threaded through a pyrrolidine[60]fullerene bearing a dibenzocrown ether [405]. However, a cathodic shift in the oxidation potential of an exTTF moiety attached by a π-conjugated bridge to pyrrolidine[60]fullerene: 158 relative to exTTF (+0.50 and +0.55 V, respectively vs. SCE) has been observed and indicates electronic communication between the two units [406] and, perhaps another related example has been reported in which the 2e− oxidation potential for exTTF is ca. 0.45 V versus SCE  [407], although such studies on a similar system with a π-conjugated phenylvinylene link showed no shift in the oxidation potential of the of the exTTF moiety [408]. A 40–60 mV cathodic shift in the first reduction potential in A-D dyads of exTTF moieties π-linked to pyrrolidine[60]fullerene also indicates electronic interaction between A and D moieties [409]. Cyclic voltammetry study of TTF [410] or exTTF [411] attached to triazolino[60]fullerene show no significant D-A interactions. Electrochemical studies of exTTF directly linked to methano[60]fullerene were reported [412]. Connecting these units similarily with longer σ-bonded linkages and A-D-A and D-A-D triads gave compounds that showed no interaction between units electrochemically [413]. Interesting studies on the effect of the linker on the interaction between two redox centers connected by acetylene moieties and Pt (two TTF moieties) [414], acetylene moieties and ruthenium (ferrocene and TTF) [415], have been reported. Electrochemical evidence for the tuning of exTTF donor acceptor complexes by aryl substituents has been presented [416]. Redox control of exTTF porphyrin donor acceptor complexes by metal ion coordination has been shown [417]. A review of the application of exTTF moieties in organic electronics has appeared [228]. Self-assembly of exTTF appended with pyridine groups by Pd(II) and Pt(II) into electroactive containers has been reported [418]. Cyclic and Osteryoung square-wave voltammetry of exTTF σ-linked to pyrene shows the same oxidation potential for exTTF oxidation in solution as for that anchored in a single-walled carbon nanotube (SWNT) by the pyrene moiety (although with broadening), but the embedded pyrene moiety shows 100 mV shift compared with that in solution [419]. Electrochemistry of TTF and exTTF moieties covalently attached via spacers to SWNTs has been measured to provide estimates of the energy required for forming the corresponding charge separated states [420]. Differential pulse voltammetry of exTTF moieties conjugated with cyanoacrylic acids provided estimates of the HOMO–LUMO gap in these compounds that after attachment to mesoporous TiO2 are used in a dye-sensitized solar cell [421]. Electrochemical studies of polythiophenes annulated with TTF obtained by electropolymerization of the monomers have been reported [422,423] as well as for oligothiophenes linked on one end to TTF and at the other end to pyrrolidino [60]fullerene [424].

VII.

THIOPHENE AND ANALOgUES

Electropolymerization of thiophene and its derivatives [425,426] is arguably the most important electrochemical behavior of thiophene because activated (“doped”) polythiophenes such as poly(ethylenedioxy)thiophene PEDOT or PEDT are the most useful organic conducting polymers [427–429]. Oligothiophenes are also of great interest as components in organic electronic devices and molecular electronics [430]. The electrochemistry of oligo- and polythiophenes has been extensively studied and recently reviewed [431]. Consequently, these areas are excluded from the present overview but are included in Chapter 41 of this volume. Anodic oxidation of thiophene and its alkyl derivatives is typically irreversible leading to electropolymerization. However, annulations of thiophene with bicyclo[2.2.2]octane moieties as in 159 results in reversible 1e− oxidation in CH2Cl2 with E1/2 = +0.79 V versus Fc/Fc+ [432]. As pointed out in Section V, the bicyclo [2.2.2]octene moieties stabilize the radical cation by inductive, σ-π conjugative and steric effects [177].

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Sulfur-, Selenium-, and Tellurium-Containing Compounds

S 159

S

160

Ar

S

Ar

161a, Ar = Ph b, Ar = p-MeOC6H4 c, Ar = p-CF3C6H4

Similarily, the dibenzothiophene derivative 160 shows a reversible 1e− oxidation in CH2Cl2 with E1/2 = +0.67 V versus Fc/Fc+ and a quasi-reversible second oxidation [433]. Compounds 161a–c exhibit a reversible 1e− oxidation as well as a reversible 1e− reduction in THF [434]. 2,5-Di(alkylthio)thiophenes also show reversible 1e− oxidations in CH2Cl2 or (CF3)2CHOH [435,436] followed by an irreversible oxidation. However, 2,2′-di(alkythio) bithiophenes show two reversible oxidations [435,436]. The effect of nitro substituents appended to thiophene and benzothiophene was evaluated by cyclic voltammetry [437]. The reduction potential of thiophenes bridged by carbonyl or 1,2-dicarbonyl groups has been recorded [438]. Bulky aryl substituents have been shown to improve the stability of the electrochemical oxidation products from dithieno-thiophenes, that is, the first oxidation waves are reversible and electropolymerization is thwarted by cyclic voltammetric studies in CH2Cl2 [439]. Reversible 1e− oxidation in acetonitrile has also been reported for alkoxysubstituted dibenzothiophene [440]. The products of further oxidation may be sulfonium salts, 5-oxides or 5,5-dioxides [440,441]. Thiophenes fused to benzenoid aromatic rings have been of much interest as semiconductors for organic field effect transistors and their electrochemistry reported in an effort to establish HOMO– LUMO energy gaps. R S

S R 162a, Ar = S

S

b, Ar = S c, Ar = Ph d, R = 2-naphthyl e, R =

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C6H5 CH3

1074

Organic Electrochemistry S

Ph

Ph

S 163 R

R

S

S

S

S R

R

164a, R = H b, R = C6H13

Thus, cyclic voltammetric studies in CH2Cl2 of 162a–e show oxidation potentials of +0.06, +0.02, +0.29, +0.22, and +0.25 V versus Fc/Fc+, respectively [442]. Compound 163 exhibits a reversible oxidation wave in PhCN with E1/2 = 0.91 V versus Fc/Fc+ [443]. Cyclic voltammetry of 164a shows irreversible oxidation owing to polymerization but if the α-sites in the thiophene moieties are blocked as in 164b, oxidation is reversible and differential pulse voltammetry shows oxidation to the radical cation followed by oxidation to the dication. Similar results occur with the isomeric system [444]. Cyclic voltammetric studies on compounds 165, R = H, C12H25 or ethylhexyl show strong reduction and weak oxidation waves from whose onset potentials LUMO and HOMO energies were estimated [445]. R S

S R 165 R΄ R

O S

S R

O

N

O

N

O

R S

S R

R΄ 166, R = C12H25, R΄= C8H17

Donor–acceptor complexes with thiophene and carborane [446], substituted anthracenes [447], or 1,3,4-thiadiazole [448] have been studied electrochemically. Electrochemical studies to assess

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1075

Sulfur-, Selenium-, and Tellurium-Containing Compounds

the role of thiophene and its position in linking donor and acceptor moieties have been studied [449]. Cyclic voltammetric studies on four substituted thiophenes annelated to anthracene [450] and three aryl substituted thiophenes annelated to naphthalene [451] have been reported. Electrosynthesis of a cyclo[9]pyrrole with three symmetrically interspaced thiophenes has appeared [452]. Electrochemical reduction of thiophene fused to a 1,3,2-dithiazolium ring afforded the corresponding radical as substantiated by EPR spectroscopy [453]. A selective electrochemical thiophene-based sensor for Ca2+ has been reported [454]. Spectroelectrochemical studies on an end-capped thiophene tetramer have been done and, on cooling, the one electron oxidized product forms a dimer whose structure was elucidated by the use of NMR spectroscopy, a spectroscopic technique rarely used in spectroelectrochemical studies [455]. The redox chemistry of oligo- and polythiophenes, polythienylene vinylenes and A-D polymers, where D is thiophene, has been reported but is beyond the scope of this review. Conjugated D-A polymers, in which the donor is benzodithiophene, have attracted much attention for use in solar cells. Monomeric 166 illustrates such a conjugated D-A motif and undergoes a reversible 1e − oxidation and two reversible 1e− reductions (the onset potentials gave a low HOMO–LUMO gap of 1.52 eV) [456]. A recent, related, nonpolymer example incorporating naphthodithiophene 167 as donor has been reported. Oxidation of 167 using cyclic voltammetry was used to estimate its HOMO energy [457]. OR΄ R

S S

R

OR΄

O

R΄ N S

167, R =

S

N

R΄

O

R΄= 2-ethylhexyl S

S

S

SiMe3 S Br S

Br S S

SiMe3

168

Cyclic voltammetric studies in CH2Cl2 of chiral [7]helicene 168 show two reversible waves at oxidation potentials much more positive than typical for alkyl oligothiophenes [458]. Spectroelectrochemistry shows that the first oxidation wave for 168 produces the corresponding radical cation that is also obtained by chemical oxidation with NOPF6. Interestingly, this radical cation is configurationally stable at room temperature. Octathio[8]circulene 169 is not soluble in organic solvents but solid state cyclic voltammetric measurements have been made on sublimed films [459] or thin films on ITO with ionic liquids [460]. Interestingly, the films show color changes depending on their oxidation state, that is, electrochromism [460].

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1076

Organic Electrochemistry S R

S

S

S

S

S X

X

S

S

R S

S

S

S R

R

170a, X = SiMe2, R = H

169

b, X = S, R = TMS

Compounds 170a and b, in which planarity of the central cyclooctatetraene can be tuned by bond length changes of the bridging X group, show oxidation at +0.45 and 0.54 V and two reduction potentials at −2.09, −2.54 and −1.79, −2.25 V, respectively, in CH2Cl2 for the oxidation and in THF for the reduction versus Fc/Fc+ [461]. The oxidation results in production of the corresponding radical cation and the reduction results in the formation of the radical anion and dianion, sequentially. Note that the less cathodic reduction potentials for 170b than 170a are ascribed to increased planarity of 170b over 170a, resulting in a lowered LUMO due to enhanced antiaromaticity of the central 8π cyclooctatetraene moiety. Finally, the electrochemistry in CH2Cl2 of imine-linked thiophenes 171 shows two consecutive 1e − oxidations to the corresponding radical cation and dication, respectively, and the first oxidations used to calculate their ionization potentials [462]. R΄

R΄

R΄ S

N

R΄

N S

S R2

O NR S (O)n

R2

171a, R΄= H; R2 = CO2Et ; R2 = CO

b, R΄= C10H21

2Et

172a, R = H, n = 0 b, R = H, n = 1 c, R = H, n = 2 d, R = Ph, n = 0 e, R = Ph, n = 1

The radical cations undergo coupling but are sufficiently stable under electrochemical conditions so that the first 1e− oxidation is quasi-reversible. In addition, the products from coupling the radical cations show reversible oxidation when deposited on the ITO electrode. Electrooxidation of tetraphenylthiophene S-oxide and reduction of thiophene-S-oxide has been reviewed recently [463,464]. The electrochemical oxidation and reduction potentials of a series of thiophenes linked to bithiophene S,S-dioxides have been reported [465].

VIII.

POLyTHIO AND SULFUR-NITROgEN UNSATURATED HETEROCyCLES

Heterocycle 172a undergoes irreversible oxidation to afford the corresponding sulfonamide 172c in CH3CN and sulfinamide 172b in the presence of Na2CO3 in modest yields. However, indirect oxidation of 172a as well as 172d using Et4NCl results in high yield of 172b and e, respectively [466]. The redox chemistry of five-membered (and larger) ring polythio and sulfur-nitrogen unsaturated heterocycles [467] has been studied primarily to obtain stable, neutral radicals that are conducting and show magnetic ordering in the solid state [468–470]. A recent review on persistent organic radicals in molecular materials has appeared [471]. Here, the focus is on their electrochemical properties in solution. The redox potentials for five–eight–membered ring heterocycles has been reviewed [467], and this review updates this previous review.

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Sulfur-, Selenium-, and Tellurium-Containing Compounds

Cyclic voltammetry of 173a in ClCH2CH2Cl shows a reversible 1e− oxidation wave with E1/2 = −0.34 V and a quasi-reversible 1e− reduction at −0.71 V versus SCE [472]. Et

E

S

S

S

S

S

S

S

S

E΄ N

S

S Et S

Sn

175a, E = E΄= S b, E = S ; E΄= Se c, E = Se; E΄= S d, E = E΄= Se

174

173a, n = 0 b, n = 1

Two reversible redox waves are reported for the hexathio-phenalenyl system 173b [473] and analogues of 173 with borate bridging two moieties and acyclic disulfides rather than cyclic disulfide moieties has been investigated [474]. Benzotrithiole undergoes reversible 1e− oxidation and 174 shows two reversible 1e− oxidations, indicating interaction between the two trithiole rings [208]. The first oxidation gives the corresponding radical cation, and the second the corresponding dication. Interestingly, the difference in energy between the singlet and the triplet dication appears to be small [475]. Cyclic voltammetric studies in CH2Cl2 of a series of 1,2,3-dithiazoles show irreversible oxidation and reduction waves [476]. However, radicals 175 undergo reversible 1e− oxidation but irreversible 1e− reduction [477]. The redox chemistry of 1,2,3-dithiazolyl radical 176 was elicited by cyclic voltammetry in CH3CN starting with the corresponding monocation salt [478–480]. The radical is stable in solution and shows two reversible oxidation waves owing to the +1/0 and +1/+2 couples for 176, X = CH, R = Me or Et; X = CCl, R = H, Me or Et; X = CF, R = Et and a reversible −1/0 couple for 176, X = CCl, R = Me or Et but an irreversible −1/0 couple for 176, X = CH, R = Me or Et; X = CF, R = Et due to S–S or S–N cleavage. O

R N

N

N S

S S

X

S

N

N

X

S

X

S

S

S S

S

N

Ph 176

177

178a, X = CH b, X = N

Compound 176, X = CCl, R = H shows anomalous behavior on reduction that is tentatively ascribed to proton tautomerization of the anion with S–S bond cleavage [479]. Interestingly, a crystallomorph of 176, X = CF, R = Et shows unusual spin crossover between the radical and the diamagnetic dimer with a 4c, 6e− S···S–S···S σ-bond [480]. Such switchable bistability may be advantageous [481]. The electrochemistry of radicals 176, X = N, R = Me or Et with two 1,2,3-dithiazoles fused to a pyrazine ring have been studied [482,483]. Cyclic voltammetry in CH3CN shows a reversible 0/+1 couple, anodically shifted from the pyridine analogues 176, X = CH, R = Me or Et by 120 mV, a reversible +1/+2 couple and, like the pyridine analogues, an irreversible reduction to the anion owing to S-S or S-N cleavage. Interestingly, 176, X = N, R = Et dimerizes forming a C–C σ-bond, which thermally isomerizes to a dimer with a 4c, 6e− S···S–S···S σ-bond [482]. Cyclic voltammetry of 177 in CH3CN reveals a reversible +1/0 couple and an irreversible −1/0 couple [484]. A series of 1,3,2-dithiazolyl radicals, including 178a and b, have been studied electrochemically. The compounds studied have one or two of these rings fused to benzene (178a), naphthalene, pyrazine (178b)

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Organic Electrochemistry

or quinoxaline [485]. Cyclic voltammetry shows a reversible 1e− oxidation (+1/0 couple) in which the 1,3,2-dithiazolyl fused to a benzene or naphthalene ring oxidizes less anodically by almost 1 V than those fused to a pyrazine or quinoxaline ring. All of the compounds studied showed an irreversible 1e− reduction (−1/0 couple) except for 179 and bis-1,3,2-dithiazolyl fused to pyrazine, which showed a reversible reduction [486]. Cyclic voltammetric studies on an analogue of 177 but with a chloro group in place of Ph show a similar behavior [487]. Cyclic voltammetric studies of 1,2,4-thiadiazinyls 180a and b in CH2Cl2 show irreversible oxidation and a reversible reduction. A prepeak, more positive than the anodic peak is due to adsorption of the reduced species [488]. O N

S

N

Ar

CF3

N N

S O 179

N

N

S

S

X4 180a, X = CI b, X = F

181

Since increasing the scan rate decreases this prepeak current, the prepeak is ascribed to adsorption of the reduced species. Similar results were obtained for 1,2,4,6-thiatriazinyls 181 but at higher concentrations dimer formation occurred resulting in a reversible oxidation process [489]. Cyclic voltammtetry of 182 in CH3CN show reversible reduction to the corresponding anion radical, which was characterized by spectroelectrochemistry (EPR spectroscopy) [490], and irreversible oxidation with the parent 182, X = CH more difficult to oxidize or reduce [491]. O(CH2)3SO2Ph N

S

X

N

N

X S

SO2 PhSO2

N OCH2OCH3

182a, X = CAr, CNMe2, CtBu, or PPh2

IX.

184

183

SULFONIUM SALTS, SULFOXIDES, SULFONES, SULFINyL, AND SULFONyL DERIVATIVES

Electrochemical reduction of triaryl, diaryl alkyl, and aryl dialkyl sulfonium salts has been studied [2] and results in C–S bond cleavage. Whether the cleavage occurs concertedly with electron transfer or stepwise via the corresponding sulfuranyl radical depends on the structure of the sulfonium salt and energy of the electron added [492]. Recent studies on indirect electrolyses mediated by cyanoaromatics show that the C–S bond cleaved in diaryl alkyl and aryl dialkylsulfonium salts depends on the sulfuranyl radical structure and bond dissociation energies [493]. The aryl radicals produced by electrochemical reduction of triaryl sulfonium salts have been grafted onto glassy carbon [494]. Electrochemical reduction of S-benzyl [495], S-allyl [496], or S-cyano [497] thiolanium or diethyl sulfonium salts produces sulfonium ylides, as shown by trapping with benzaldehyde, which then rearrange. The mechanism for deprotonation in these reactions is unclear, but electrochemical base generation has been effectively used to provide the anion of DMSO. Such generation from DMSO and KClO4 has been reported and subsequent reaction of the anion studied [498,499]. Electrochemical reduction of sulfones has been extensively studied [2]. For aryl alkyl sulfones, the corresponding anion radical is formed that preferentially undergoes alkyl C–S bond cleavage. However, electron-withdrawing substituents on the aromatic ring favor aryl C–S bond cleavage  [500,501].

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Sulfur-, Selenium-, and Tellurium-Containing Compounds

Low-temperature electrochemistry and fast scan voltammetry in THF of sulfone 183, which undergoes alkyl C–S cleavage, facilitated chronoamperometric and voltammetric analysis [502]. Cyclic voltammetric studies on the reduction of 2-methyl-thioxanthen-9-one, S-oxide and S,S-dioxide have been reported and studies on their radical anions [503]. Aryl polysulfones undergo complex behavior, including C–S cleavage, coupling, and self-alkylation. The electrochemical results have been correlated with DFT calculations [504]. With hexakis(alkylsulfonyl)benzenes, the electrochemistry is complicated by the presence of conformers of both the neutral and anion radical forms that can be modeled by a four-membered square scheme [505]. The anion of allyl phenyl sulfone is formed by electrogenerated base. Reprotonation at the γ-position results in isomerization to PhSO2CH=CHCH3, which undergoes nucleophilic addition with allyl phenyl sulfone anion. Such nucleophilic addition also occurs with this anion generated in the presence of phenyl or tolyl vinyl sulfone. Since such nucleophilic additions generate a carbanion that regenerates the initial base, the electrogenerated base serves as a catalyst [506]. The yields are high for these reactions and even the eight-membered ring 184 is formed in 41% yield by this reaction of allyl phenyl sulfone with divinyl sulfone. Alternatively, aryl vinyl sulfones undergo [2+2] cycloaddition on electrochemical reduction via an ECE mechanism featuring the initial formation of the corresponding radical anion [507]. This reaction is also catalytic because the product radical anion transfers an electron to the vinyl sulfone. This electrochemical [2+2] cycloaddition has been extended to the reactions of vinyl sulfones with α, β-unsaturated ketones although in modest yields [508]. Analogues of 182 with X = C(2-thienyl) and other analogues show irreversible reduction and oxidation with no evidence for electropolymerization [509]. Anodic oxidation of N-monoalkyl p-toluenesulfinamides [RNHS(O)Ar] under controlled current conditions in MeOH affords the corresponding sulfonamides (S-oxidation) in good yields, whereas such oxidation of N,N-disubstituted and N-phenyl p-toluenesulfinamides results in N–S cleavage to produce the corresponding amine and methyl p-toluenesulfinate [510]. The difference in products is ascribed to preferential electron transfer from sulfur in the first case (S-oxidation) and from nitrogen in the second case. However, in the presence of base (KHCO3) α-deprotonation follows N-oxidation of chiral N-arylsulfinyl piperidine to afford α-methoxy-N-sulfinyl piperidine 185 in good yield as mixture of diasteromers as well as small amounts of N–S cleavage products [511]. R

Ar

O

R

S N N

N

TsNHCH2CH2 NR

OMe N N R

185

186a, R = Ts b, R = H

2 R 187a, R = Ts b, R = H

Electrochemical studies on sulfonamides have been reviewed [512,513] and the mechanism for electrochemical reduction of aryl sulfonyl phthalimides explored [514]. Of particular importance is the electrochemical reductive cleavage of the SO2–N bond. A detailed study on the potentiostatic cleavage, in CH3CN using carbon electrodes, of the sulfonamide moieties in 186a to give 186b in 80% yield and 55% faradaic efficiency has recently been reported, and the effects of supporting electrolyte, concentration of substrate and proton donor elucidated [515]. Good yield of reductive detosylation of a variety of arenesulfonamides was achieved at constant current in DMF in an undivided cell, Pt cathode and naphthalene mediator [516]. Selective detosylation of 187a to 187b was also achieved in good yield. Electrochemical reduction of N-phenylsulfonyl N-substituted (R)-phenylglycinol in MeCN with Bu4N+HSO4− at constant current at a Hg cathode has been reported [517]. The mechanism for ArSO2–NRR′ electroreduction is shown in Equation 27.17. Evidence for this mechanism was obtained by the the observation of β-elimination from the intervening RR′N• species.

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1080

Organic Electrochemistry +e– ArSO2NRR΄

ArSO2NRR΄ +e–

RR΄N

RR΄N

+ ArSO2–

(27.17)

H+

RR΄N–

RR’NH

The mechanism for electrochemical reduction of sulfonyl fluorides [518] and sulfonyl chlorides [519] using dissociative electron-transfer theory has been reported.

X. SELENIUM AND TELLURIUM COMPOUNDS Electrochemical studies of organoselenium and organotellurium compounds have been recently reviewed in an article discussing chemical and electrochemical oxidations and reductions of the compounds [520]. Cyclic voltammetric studies of 188, Fc3P = Se (Fc = ferrocene) and (tBu)2(2PhC6H4)P = Se in CH2Cl2 show an irreversible oxidation by an EE mechanism to give the corresponding product with an Se–Se bond (this product with an intramolecular Se–Se bond obtained from 188, R = t-Bu by chemical oxidation, was unequivocally characterized by x-ray crystallographic analysis) [521–523]. Se PR2 R S

R Fe

Se

N O

PR2

CO2CH2(pMeOC6H4) 189a, R = Br b, R = H

188

Electrochemical reduction of Ph2Se2 produces PhSe−, which debrominates 189a to 189b. Since Ph2Se2 is reformed in this reaction, it is used as a catalytic mediator [524]. Ph2Te2 may be used similarly as an electrochemical reduction mediator. Another sequence initiated by electrochemical reduction of Ph2Se2 occurs in the reaction shown in Equation 27.18. Selective reduction of Ph2Se2 to generate RCH 2 =CH 2 + CF2 BrX + Ph 2Se 2 → XCF2CH 2CH(SePh)R

(27.18)

PhSe− is followed by an SET mechanism to yield the product [525,526]. Bromine reacts with Ph2Se2 to produce PhSeBr, which is a useful electrophile for effecting alkoxyselenylation. Furthermore, the addition products so obtained can be oxidatively cleaved resulting in the transformation shown in Equation 27.19. OMe Ar

CO2R

Ar

CO2R

(27.19)

(or corresponding lactone)

This entire sequence can be effected electrochemically using 10 mol% of Ph2Se2 and Et4NBr as the redox catalyst in MeOH with a catalytic amount of H2SO4 [527–531]. A related electrochemical process can be effected with Ph2Se2, Et3N · 3HF and internal aliphatic alkenes (and alkynes) [526,532]. Here, Ph2Se2 is electrooxidized in CH2Cl2 and reacts with F− to form PhSeF that fluorosulfenylates the alkene (or alkyne). Further electrochemical oxidation effects elimination to afford the allylic fluoride. With electron-deficient alkenes, this procedure does not work but oxidation of Ph2Se2

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1081

Sulfur-, Selenium-, and Tellurium-Containing Compounds

with Et3N · 5HF in CH3NO2 again generates PhSeF that fluoroselenylates electron-deficient alkenes (α,β-unsaturated esters, acids, amides, and phosphonates) whose regiochemistry depends on substituents [533]. Electrophilic selenium produced on electrochemical oxidation of Ph2Se2 in MeOH readily reacts with enols, resulting in the selenium-catalyzed overall transformation of ArCOCH3 to ArCOCH(OMe)2 [534]. The quasireversible oxidation potentials for 2,7-dimethoxynaphthalene-1,8-dichalcogenides (S, Se and Te) have been reported [535]. As with thioethers, anodic oxidation of seleno- and telluroethers involves an EC mechanism [520]. However, as outlined for dithioether 33, diselena- and ditelluraethers 190a and b, respectively, undergo two reversible one-electron oxidations [536,537] with potential inversion [84,538].

M

X M

X΄

Se

190a, X = X’ = S e; M = M’ = C H 2 b, X = X’ = Te ; M=CH2 c, X = S , X’ = S e M=CH2 d, X = S , X’ = Te ; M=CH2

Se

Se

Se

191

Se Se

192

Se

Se Se

193

Interestingly, the mixed dichalcogenaethers 190c and d undergo irreversible 1e− oxidation and a reduction peak associated with the oxidation. These results are explained by rapid dimerization of the radical cations of 190c and d to give dications with an Se+–Se+ and Te+–Te+ bond, respectively. Thus, 190a and b form dications with intramolecular Se+–Se+ and Te+–Te+ bonds respectively; whereas 190c and d form dications with intermolecular Se+–Se+ and Te+–Te+ bonds rather than intramolecular Se+–S+ and Te+–S+ bonds, respectively. These results are ascribed to thermodynamic control of dication formation because the relative bond strengths are Te+–Te+ > Se+–Se+ > S+–S+ and Te+–Te+ > Te+–S+ and Se+–Se+ > Se+–S+ based on calculations [539]. Polyselenides 191–193 showed reversible oxidations at potentials that are substantially less anodic than the irreversible oxidations of 1-methylselenylnaphthalene or diphenyl selenide [540]. The facilitated oxidations are ascribed to lone pairlone pair destabilization in the neutral compounds and bond formation in the oxidized compounds. 2+ XPh TePh

XPh 194a, X = S b, X = Se

XPh TePh

Se

Se

Se

Se

XPh 195a, X = S b, X = Se

196

Cyclic voltammetric studies in acetonitrile of 194a and b show pseudoreversible oxidations with oxidation potentials 234 and 277 mV less anodic, respectively, that for Ph2Te which undergoes irreversible oxidation [541]. Since chemical oxidation (NOBF4) give dications 195a and b, respectively, whose structure was unequivocally established by X-ray crystallographic studies, it is surmised that these are formed on electrochemical oxidation as well. Such 1,2-di-chalcogenadications and related species have been reviewed [542,543], and hypervalent bonding as displayed in 195 has also

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Organic Electrochemistry

been reviewed [544,545]. Tetraselenaether 196 shows two quasi-reversible 1e− oxidations separated by 0.9 V first to the radical cation and then to the dication [546]. Unlike 190a, there is no potential inversion. Compounds 190; X = Se, M = SiMe2 or SnMe2 and X = Te, M = SiMe2 or Si(SiMe3)2 have been studied electrochemically [84]. The selenium compounds show irreversible oxidations owing to C–Si or C–Sn bond cleavage as shown by chemical oxidation [547]. However, the tellurium compounds show reversible oxidations [84]. The chemical steps following oxidation of selenides may be nucleophilic attack, loss of an α-proton, bond cleavage (C–Se), dimerization or disproportionation. Anodic oxidation of diarylselenides in the presence of water affords the corresponding selenoxide or its hydrate [548]. However, disproportionation of the intervening radical can occur followed by electrophilic aromatic substitution leading to PhSeC6H4Se+Ph2 or Ph2Se+Ar in the presence of excess benzene or toluene in the case of Ph2Se+ • [549]. Although PhSeC6H4Se+Ph2 could arise via dimerization of Ph2Se+ •, the disproportionation mechanism is favored by electrochemical analysis. The oxidation potentials of diaryl selenides and tellurides have been recently reviewed and discussed [520]. Further electrochemical determinations of the oxidation potentials of such compounds has been aimed at correlating these values with catalytic antioxidant activity [550] and ability of these compounds to serve as biological redox modulators or sensitizers [551,552] in the treatment of diseases associated with oxidative stress [553]. For example, compounds 197 show a quasi-reversible reduction due to the quinone moiety and an irreversible oxidation due to the chalcogen on glassy carbon [554]. The oxidation potentials are rendered more positive than the corresponding ArXAr, X = Se, Te owing to the electron-withdrawing quinone moiety. Electrochemical oxidation of arylmethyl selenides typically results in an ECE mechanism with α-proton loss as the chemical step. This loss of an α-proton can be suppressed by the addition of acid, in which case bimolecular reactions of the intervening radical cation occur [555]. Nevertheless, the selenium analogue 198 of Wurster’s blue undergoes selenoxide formation of both selenium moieties at different oxidation potentials [556]. O

SeMe XAr O 197a, X = Se; Ar = p-MeOC6H4 b, X = Te; Ar = Ph c, X = Te; Ar = p-MeOC6H4

PhSeCXX’CO2Et SeMe 198

199a, X = X’= H; b, X = F; X’= H; c, X = X’= F;

Perhaps greater charge delocalization in this case mitigates α-deprotonation. In “superdry” CH3CN, 198 reversibly oxidizes to its radical cation that reversibly dimerizes. α-Proton loss is favored by electron-withdrawing groups (CN, CO2Et, CONH2) appended to the α-carbon. Consequently, anodic oxidation of 199a in CH2Cl2 with Et3NHF leads to sequential α-mono- and α,α-difluoro compounds 199b and c [557,558]. Electrochemical oxidation of aryl alkyl selenides can also lead to Se–C cleavage of the radical cation, particularly with substituted alkyl groups that form more stable carbocations on cleavage, yielding ArSe•, which leads to ArSeSeAr. However, electrochemical studies suggest that the oxidation of ArSeCH2SiMe3 results in ArSeSeAr by dimerization of the intermediary radical cation followed by Se–C heterolysis [559]. As outlined earlier, thioacetals and ketals undergo electrochemical oxidation via an EC mechanism with α-C–S cleavage of the radical cation owing to the stabilization of the resulting carbocation by oxygen. Thus seleno- and telluroglycosides serve as glycosyl donors on anodic oxidation. Electrochemical oxidation of seleno- [142,144,145] and telluroglycosides [144,145] in the presence of partially protected

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Sulfur-, Selenium-, and Tellurium-Containing Compounds

monosaccharides results in disaccharide formation. Since the oxidation potentials for chalcogenoglycosides follow the order S > Se > Te (and for the same chalcogen correlates with Hammett σ+ for p-substituted phenyl chalcogenoglycosides [142]), selective electrochemical activation is possible. OBn O BnO BnO

OBn O

BnO

OBn O

+

SePh

STol

BnO

OBn

BnO 200

+1.7V

BnO

O BnO BnO BnO

201

(27.20) STol OBn

202

In this way, 200 was coupled with 201 at a potential of +1.7 V versus SCE in acetonitrile (a solvent that favors β-glycoside formation) to give 202 as shown in Equation 27.20. Further coupling of 202 with 203 at a potential of +2.0 V verusus SCE resulted in the expected trisaccharide [145]. O

OH O

O O

203

RO

X

OR

RO

Y

OR

X

O

Y 204a, X = Y = Se b, X = Se; Y = SeO c, X = Se ;Y = S d, X = SeO; Y = S

205

Although trisaccharide synthesis could be achieved, an interesting problem was uncovered in the first step. That is, electrochemical oxidation of 200 produces PhSeSePh. Furthermore, the oxidation potential of PhSeSePh is 1.45 V versus SCE under these conditions and its oxidation product activates thioglycosides, even though they are not directly oxidized. This side reaction results in a low yield (35%) of 202. Redox catalysis of phenylselenoglycoside oxidation has been reported [560]. The electrochemical mediator was tris(p-bromophenyl) amine. Interestingly the corresponding triaryl aminium radical cation was not a good enough oxidant to effect the desired oxidation, but in the presence of 2,6-lutidine, oxidation occurred. Electrochemical generation of H2Se using a selenium cathode was used in the synthesis of 2,4,6-triphenylselenopyran [561]. The electrochemistry of extended chalcogenapyrans has attracted interest [562]. Chalcoxanthylium dyes are of interest in solar cells and electrochemical studies on them [563] and related dyes reported. Selenium and tellurium analogues of thianthrene 71a have been studied electrochemically. Selenanthrene 204a undergoes two irreversible oxidations in acetonitrile under conventional conditions and controlled potential electrolysis gives the corresponding oxide 204b, but the first oxidation is reversible at high scan rates [564]. Thioselenanthrene 204c undergoes two oxidation processes in acetonitrile. Controlled potential electrolysis shows that the first oxidation is a two-electron oxidation, and spectroscopic studies provide evidence for the formation of selenoxide 204d. Detailed electrochemical studies support a disproportionation mechanism, as discussed earlier, for the reaction of nucleophiles with thianthrene radical cation, in which the radical cation oxidizes the hydroxylated radical cation [565]. The order of reactivity of the radicals with water is selenanthrene > thioselenanthrene > thianthrene. This reactivity sequence was extended to phenoxathiin 204, X = O, Y = S and phenoxaselenin 204, X = O, Y = Se radical cations in which the former is comparable to thianthrene and the latter to selenanthrene radical cations: (Se, Se)+• ~ (Se, O)+• < (Se, S)+• ≫ (S, O)+• ~ (S, S)+• [565]. In a similar way, alkoxylation of chalcogenanthrenes, as in 205, has been shown to lower their oxidation potential and stabilize the corresponding radical cations [566]. Like selenanthrene, dibenzo[c,e]-1,2-diselenin 206 shows two oxidation peaks on cyclic voltammetry and

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Organic Electrochemistry

the first oxidation is a 2e− process that yields the corresponding SeSe(O) species [567]. The initially formed radical cation is more reactive toward water than the isomeric 204a radical cation. A reduction step believed to be due to Se–Se cleavage is also observed as well as an associated oxidation. R

206

Y Y

X

R

Se-Se

X

Y

R R

207a, X = Y = Se; R = H b, X = Y = Te; R = H c, X = S; Y = Se; R = H d, X = Y = Se; R = Me e, X = Y = Se; R = CnH2n+1

Y

X

X

X

X

Y Y

208a, X = Se; Y = O b, X = S, Y = O

The selenium and tellurium analogues of tetrathiafulvalene, 207a and b, have been studied electrochemically as well as mixed chalogen species, for example, 207c and reviewed previously [1,230,568,569]. Two reversible 1e− oxidations are observed and, interestingly, the first oxidation potential follows the order 82 lower than 207b lower than 207a and ΔE for the first and second oxidation potentials follow the order 207b < 207a < 82. The first oxidation potentials show that, despite the lower ionization energy of Se and Te compared with S, the π-delocalization in the radical cation dominates Eo′. Better overlap of the sulfur 3p orbitals than Se 4p or the Te 5p orbitals with carbon 2p orbitals accounts for this difference in delocalization. It is also notable that the ΔE values that reflect the Coulombic repulsion in the dications are lower for Te than Se than S. These compounds are of interest because of the conductivity and magnetic properties of their salts, as already discussed for the sulfur analogues. Indeed, the first molecule-based superconductors were derivatives of 207a [212,213]. The Bechgaard salts are synthesized by electro- or chemical oxidation. A recent interesting example involves the electrooxidation of 207d, in an ionic liquid as electrolyte to prepare (207d)2 NbF6 [570]. The effect of substituents on the oxidation potentials of 207a and b has been reviewed and obey the Hammett equation [1,230,568,569]. The standard redox potentials for 207a and some of its derivatives have been calculated using DFT and a polarized continuum model for the solvent [571]. Owing to the importance of ethylene-dioxytetrathiafulvalene and ethylene dithioxytetrathiafulvalene, the corresponding selenium compounds and analogues have been studied as well. Thus, 208a undergoes two 1e− reversible oxidations at lower potentials due to the electron-donating substituents than the parent 207a but at slightly more positive potentials than the corresponding 208b [572]. Galvanostatic oxidation was used for synthesis of the corresponding cation radical salts [573]. Alkylthio substituents appended to tetraselenafulvalene render the oxidation potentials more positive in 207e, but the length of the alkyl chain (C1–C15) has little effect [574]. Cyclic voltammetric studies of 208, X, Y = S, Se and analogues (mixed S,Se compounds) in benzonitrile show two reversible 1e− oxidations with the oxidation potentials, particularly E01 , more anodic than the corresponding parent [575]. The electrochemistry of 209 and analogues was reported [576,577] and, as expected, the methyl groups lower the oxidation potential and the sulfur substituents raise it. However, the corresponding cation radical perchorate shows extraordinary conductivity (1500 S/cm) at room temperature. S Me Me

Se

Se

S

X

Se

Se

S

X

209

R R

S S S

R R

210a, X = Se; R = H b, X = Se; R = CO2Me c, X = Se; R , R = S(CH2)2S

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Sulfur-, Selenium-, and Tellurium-Containing Compounds

π-Extended (dendralene) systems 210a–c have been reported and studied by cyclic and Osteryoung square-wave voltammetry in CH2Cl2 [578,579]. Three waves are reported. The interesting 1,3-ditelluretane π-extended system 211 as a mixture of E/Z isomers was studied using cyclic voltammetry in CH2Cl2 [580]. A reversible oxidation followed by three irreversible oxidation peaks was reported. Cyclic voltammetric studies in benzonitrile of bitetetraselenafulvalene 212 show two oxidation peaks each apparently involving 2e− oxidations with the second oxidation irreversible [581]. Electrocrystallization gave conducting salts.

S

S Se

Te S

Se

Se

Se

R

Se

Se

X

Y

S Te

Se

211

Se

Se

212

213a, X = Y = O; R = H b, X = O; Y = S; R = H c, X = O; Y = S; R = Me d, X = Y = S; R = H

Owing to the great interest in polythiophenes and especially PEDOT [poly(ethylenedioxy)thiophene], the electrochemistry and electropolymerization of selenophene and its derivatives has been studied [582–585]. Electrochemical or chemical oxidation of 3-alkylselenophenes provides a regiorandom polymer. However, chemical coupling reactions provide regioregular poly(3-hexyl)selenophene [586]. From the onset of the oxidation and reduction peaks obtained by cyclic voltammetry of the polymer, the band gap was determined (which was 0.3 eV larger than the optical band gap) and is smaller than the polythiophene analogues. Selenophene derivative 213a shows an irreversible oxidation at 1.18 V in CH3CN versus SCE, which is less positive than that for the thiophene analogues (1.44 V) [587,588]. Repetitive scans lead to a deep blue polymer deposited on the electrode surface. Oxidative polymerization could also be achieved by chemical oxidation with FeCl3. Other selenophenes 213b–d behave similarily, that is, they have lower irreversible oxidation potentials than the corresponding thiophenes (by ca. 100 mV) and electropolymerize [589,590]. The band gaps of these polymers are determined spectroelectrochemically and, despite the changes in the peripheral atoms, are all approximately 1.4 eV. The data and calculations for these polymers argue for greater planarity in the backbone of polyselenophenes than polythiophenes due to increasing the extended π-conjugation in the selenium polymers as a result of greater “quinoid” character. OMe

MeO

Te Te 214

R

R

R

R

215, R = C6H13

Cyclic voltammetric studies of 3,4-dimethoxytellurophene 214 show two oxidation peaks at potentials well below the corresponding selenophene and thiophene analogues [591]. However, electropolymerization of 214 proved problematic. Consequently, a polytellurophene derivative was prepared by Pd-catalyzed coupling reactions and its HOMO level was determined from the onset of its oxidation determined electrochemically [592]. However, its degree of polymerization is less than that achieved with thiophenes [593]. Tellurophene derivative 215 and its Br2 adduct have been studied by cyclic voltammetry [594]. Two reversible 1e− oxidations are observed for 215, and its Br2 adduct shows irreversible oxidation and reduction. The HOMO energy for the Br2 adduct is lower than that for 215 based on the oxidation onsets. Reversible electrochemical oxidation in water of a watersoluble tellurophene has been reported [595]. Cyclic voltammetric studies were used to determine frontier molecular orbital energies for a 2,5-diaryltellurophene and its corresponding tellurium

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Organic Electrochemistry

dibromide [594]. Electropolymerization of 3-hexyltellurophene has been studied and its regioregular polymer, obtained by chemical coupling, has been examined electrochemically to assess its HOMO–LUMO gap [596]. Cyclic voltammetric studies of benzodichalcogenophenes 216a–c show irreversible oxidation peaks with the tellurium analogue oxidizing at much lower potential (+0.48 V vs. Fc/Fc+) than the selenium (+0.89 V) or sulfur analogues (+0.95 V) [597] (a second oxidation peak was also observed for the tellurium compound 216c). Se

X

R

R

Se X

R

R

216a, X = S; R = H b, X = Se; R = H c, X = Te; R = H d, X = Se; R = Ph

Se

Se

Se

Se 217a, R = H b, R = Ph

218

The onset of their anodic peak was used for estimating HOMO levels in these compounds. Similar studies on the directly fused diselenophenes 217a and b have been reported [598], and their irreversible oxidation (+1.06 and +0.86 V vs. Fc/Fc+, respectively) occur at higher potential than that of 218 (+0.41 V) or 216d (+0.80 V). Interestingly, high-performance organic field effect transistors were fabricated with 217b [598]. Two reversible 1e− reductions of 219a–c are shown by cyclic voltammetric studies in acetonitrile [588,599]. R O

R N

219a, R = CN b, E = CO2Et c, E = CHO

N

N X

X Se

N

N

NH

R

Se 220

Y

Y Cl

221a, X = S; Y = Se; R = Et b, X = Se; Y = S; R = Et c, X = Y= Se; R = Et d, X = Y= Se; R = Me

Cyclic voltammetric studies on the [8]circulene (see 169) in which thiophene and selenophene rings alternate have been reported, and its oxidation potential is substantially less anodic than that for the all thiophene 169 [459]. The electrochemistry of selenium–nitrogen heterocycles has attracted attention. Cyclic voltammetric studies on 220 show a prepeak at +1.24 V in acetonitrile and a well-defined irreversible oxidation at +1.44 V [600]. Controlled potential electrolysis provided the corresponding SeO compound. As already pointed out, selenium radicals 175b–d undergo reversible 0/+1 but irreversible 0/−1 couples [477]. The redox chemistry of 1,2,3-dithiazolyl radicals 176 has been discussed and the selenium analogues 221 have also been studied electrochemically [601]. Starting with the corresponding cations, three reversible waves owing to the 0/+1, 0/−1, and +1/+2 couples are observed as is the case for the all sulfur analogue 176, X = Cl, R = Et. For 221c, two additional more positive reversible waves are found as well. There is an anodic shift in the 0/+1 couple with increasing selenium content in the series 176, X = Cl; R = Et, 221a–c. Electrochemical reduction was found to be the best way for converting nonafluorobutanesulfonate (−ONf) salt 222 into the corresponding radical 221d [602]. The redox chemistry of GaCl4− salts of 223a–c was studied in acetonitrile using cyclic voltammetry [603]. A reversible +1/0 wave is observed for 223a and b with comparable E1/2 (0.260 and 0.283 V vs. SCE, respectively). With 223c, a reversible +1/0 couple is not observed, only a weak peak at more anodic potentials than 223a and b is seen.

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1087

Sulfur-, Selenium-, and Tellurium-Containing Compounds Me N

N

N

N Se+

Se

E

E S+

Se

Se

N

–ONf

NMe+ 224a, E = S b, E = Se c, E = Te

223a, E = S b, E = Se c, E = Te

Cl 222

EPR spectroscopy of the radicals produced on reduction of 223a and b but not c support their structures. It is suggested that the radical produced by reduction of 223c is dimeric in solution. Compounds 223a and b also show an irreversible 0/−1 couple (E1/2 −0.96 and −0.90 V, respectively). Similar behavior was observed for the redox chemistry of 224a–c [604]. That is, cyclic voltammetry in CH3CN shows a reversible +1/0 wave for 224a and b (E1/2 −0.351 and −0.308 V vs. SCE, respectively) and a much more anodic, although reversible, wave for 224c. R' –

R2

N

F N

E N

S N

X R3

N R4

225a, E = S b, E = Se c, E = Te

N

F

226a, R1 = R2 = R3 = R4 = H; X = S b, R1 = R2 = R3 = R4 = H; X = Se c, R1 = R2 = R3 = R4 = F; X = S d, R1 = R2 = R3 = R4 = F; X = Se e, R1 = R2 = R4 = F; R3= CF3; X = S f, R1 = R2 = H; R3 = R4 = F; X = S g, R1 = R3 = R4 = F; R2 = CH3; X = Se h, R1 = R2 = R4 = H; R3 = CF3; X = Se

F

F F

F

227

EPR spectra of radical anions 225a and b obtained on chemical reduction of 226a and b validated their structures. Radical 225 was not observed, but its dimer was suggested to persist in solution. The redox potentials of a series of sulfur and selenium heterocycles 226a–h, 227 in acetonitrile by cyclic voltammetry has been reported [605]. The first reduction peak is reversible and the EPR spectra of the radical anions produced electrochemically have been measured. Several of these heterocycles 226b–e, g, 227 show a second irreversible reduction at more negative potentials. An irreversible oxidation is also observed. The analogous selenium compounds show more positive reduction and more negative oxidation potentials than the corresponding sulfur compounds in this series. The reduction potentials for these compounds linearly correlated with the calculated EA values (with separate lines for the sulfur 226a, c, e, f and selenium 226b, d, g, h compounds).

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4. Spataru, N.; Sarada, B. V.; Popa, E.; Tryk, K. D. A.; Fujishima, A. Anal. Chem. 2001, 73, 514–519. 5. Chailapakul, O.; Aksharanandana, P.; Frelink, T.; Einaga, Y.; Fujishima, A. Sens. Actuators B Chem. 2001, 80, 193–201. 6. Enache, T. A.; Oliveira-Brett, A. M. Bioelectrochemistry 2011, 81, 46–52. 7. McClintock, C.; Kertesz, V.; Hettich, R. L. Anal. Chem. 2008, 80, 3304–3317. 8. Chiku, M.; Ivandini, T. A.; Kamiya, A.; Fujishima, A.; Einaga Y. J. Electroanal. Chem. 2008, 612, 201–207. 9. Zhou, M.; Ding, J.; Guo, L.-P.; Shang, Q. Anal. Chem. 2007, 79, 5328–5335. 10. Zhao, Y. D.; Zhang, W.-D.; Chenand, H.; Luo, M. Q. Sens. Actuators B Chem. 2003, 92, 279–285. 11. Zhang, F.; Wen, M.; Cheng, M.; Liu, D.; Zhu, A.; Tian, Y. Chem. Eur. J. 2010, 16, 11115–11120. 12. Nekrassova, O.; Allen, G. D.; Lawrence, N. S.; Jiang, L.; Jones, T. G. J.; Compton, R. G. Electroanalysis 2002, 14, 1464–1469. 13. Guo, Z.-N.; Yao, H.-Q.; Liu, W.-Y. Electroanalysis 2005, 17, 619–624. 14. Ardakani, M. M.; Rahimi, P.; Karami, P. E.; Zare, H. R.; Naemi, H. Sens. Actuators B Chem. 2007, 123, 763–768. 15. Shaidarova, L. G.; Ziganshina, S. A.; Budnikov, G. K. J. Anal. Chem. 2003, 58, 577–582. 16. Prasad, K. S.; Muthutanian, G.; Zen, J.-M. Electroanalysis 2008, 20, 1167–1174. 17. Guppi, M. A.; Paez, M. A.; Costamagna, J. A.; Cardenas-Jiron, G.; Bedioui, F.; Zagal, J. H. J. Electroanal. Chem. 2005, 580, 50–56. 18. Pereira-Rodriguez, N.; Cofre, R.; Zagal, J. H.; Bedioui, F. Bioelectrochemistry 2007, 70, 147–154. 19. Griveau, S.; Albin, V.; Pauporte, T.; Zagal, J. H.; Bedioui, F. J. Mater. Chem. 2002, 12, 225–232. 20. Lawrence, N. S.; Davis, J.; Compton, R. G.; Jiang, L.; Jones, T. G. J.; Davis, S. N. Analyst 2000, 125, 661–663. 21. Nielsen, L. T.; Vase, K. H.; Dong, M.; Besenbacher, F.; Pedersen, S. U.; Daasbjerg, K. J. Am. Chem. Soc. 2007, 129, 1888–1889. 22. Fotouhi, L.; Hekmatshoar, R.; Heravi, M. M.; Sadjadi, S.; Rasmi, V. Tetrahedron Lett. 2008, 49, 6628–6630. 23. Chou, J.-H.; Rauchfuss, T. B.; Szczepura, L. F. J. Am. Chem. Soc. 1998, 120, 1805–1811. 24. Breitzer, J. G.; Smirnov, A. I.; Szczepura, L. F.; Wilson, S. R.; Rauchfuss, T. B. Inorg. Chem. 2001, 40, 1421–1429. 25. Torii, S. Electroorganic Reduction Synthesis, Vol. 1, Chapter 6. Tokyo, Japan: Wiley-VCH Kodansha, 2006. 26. Ralph, T. R.; Hitchman, M. L.; Millington, J. P.; Walsh, F. C. J. Electroanal. Chem. 1994, 375, 17–27. 27. Ralph, T. R.; Hitchman, M. L.; Millington, J. P.; Walsh, F. C. J. Electroanal. Chem. 2005, 583, 260–272. 28. Ralph, T. R.; Hitchman, M. L.; Millington, J. P.; Walsh, F. C. J. Electroanal.Chem. 2006, 587, 31–41. 29. Saez, A.; Sánchez- Sánchez, C. M.; Solla-Gullón, J.; Expósito, E.; Montiel, V. J. Electrochem. Soc. 2009, 156, E154–E160. 30. Brevet, D.; Mugnier, Y.; Samreth, S.; Dellis, P. Carbohydr. Res. 2003, 338, 1543–1552. 31. Peng, Z.; Holm, A. H.; Nielsen, L. T.; Pedersen, S. U.; Daasbjerg, K. Chem. Mater. 2008, 20, 6068–6075. 32. See footnote 34 in Reference 31. 33. Maran, F.; Wayner, D. D. M.; Workentin, M. S. Adv. Phys. Org. Chem. 2001, 36, 85–166. 34. Antonella, S.; Maran, F. Chem. Soc. Rev. 2005, 34, 418–428. 35. Savéant, J.-M. Adv. Phys. Org. Chem. 2000, 35, 177. 36. Daasbjerg, K.; Jensen, H.; Benassi, R.; Taddei, F.; Antonello, S.; Gennano, A.; Maran, F. J. Am. Chem. Soc. 1999, 121, 1750–1751. 37. Antonello, S.; Benassi, R.; Gavioli, G.; Taddei, F.; Maran, F. J. Am. Chem. Soc. 2002, 124, 7529–7538. 38. Antonello, S.; Daasbjerg, K.; Jensen, H.; Taddei, F.; Maran, F. J. Am. Chem. Soc. 2003, 125, 14905–14916. 39. Asmus, K. D. In Sulfur-Centered Reactive Intermediates in Chemistry and Biology (C. Chatgilialoglu, K. D. Asmus, eds.). New York: Plenum, 1990, pp. 155–172. 40. Meneses, A. B.; Antonello, S.; Arevalo, M. C.; Concepcion, C.; Sharma, J.; Wallette, A. N.; Workentin, M. S.; Maran, F. Chem. Eur. J. 2007, 13, 7983–7995. 41. Gavrilov, A. B.; Kovalev, I. P.; Skotheim, T. A. Proc. Electrochem. Soc. 1997, 96–17, 204–212. 42. Tani, H.; Yoshitoku, D.; Inamasu, T.; Ono, N. Electrochem. 2003, 71, 1055–1057. 43. Inamasu, T.; Yoshitoku, D.; Sumi-otorii, Y.; Tani, H.; Ono, N. J. Electrochem. Soc. 2003, 150, A128–A132. 44. Kiya, Y.; Hatozaki, O.; Oyama, N.; Abruña, H. D. J. Phys. Chem. C 2007, 111, 13129–13136. 45. Earlier studies had shown that the redox chemistry of 1,3,4-thiadiazole- 2, 5-dithiol was accelerated in a composite electrode with polyaniline: Oyama, N.; Tatsuma, T.; Sato, T.; Sotomura, T. Nature 1995, 373, 598–600.

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Sulfur-, Selenium-, and Tellurium-Containing Compounds 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. 61. 62. 63. 64. 65. 66. 67. 68. 69. 70. 71. 72. 73. 74. 75. 76. 77. 78. 79. 80. 81. 82. 83. 84. 85. 86.

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Sulfur-, Selenium-, and Tellurium-Containing Compounds 139. 140. 141. 142. 143. 144. 145. 146. 147. 148. 149. 150. 151. 152. 153. 154. 155. 156. 157. 158. 159. 160. 161. 162. 163. 164. 165. 166. 167. 168. 169. 170. 171. 172. 173. 174. 175. 176. 177. 178. 179. 180. 181. 182. 183. 184. 185. 186. 187. 188. 189. 190.

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28

Aliphatic Nitrogen–Containing Compounds Amines, Amino Alcohols, and Amino Acids Osamu Onomura

CONTENTS I. II.

Introduction ........................................................................................................................ 1103 Aliphatic Amines ............................................................................................................... 1104 A. Primary Amines ......................................................................................................... 1104 B. Secondary and Tertiary Amines ................................................................................ 1104 III. N-Protected Amines ........................................................................................................... 1105 A. Regioselectivity .......................................................................................................... 1105 B. Electrochemical Cyanation ........................................................................................ 1106 C. Electrochemical α,β-Functionalization ...................................................................... 1107 D. Preparation of Imides ................................................................................................. 1107 IV. N-Protected α-Amino Acids............................................................................................... 1108 A. Chemoselectivity ........................................................................................................ 1108 B. Enantioselectivity ....................................................................................................... 1108 V. N-Protected β-Amino Alcohols...........................................................................................1110 VI. N-Protected α-Allyl or Benzyl Amines...............................................................................1111 VII. Intramolecular Anodic Coupling of N-Protected Amines ..................................................1111 A. Carbon–Carbon Bond–Forming Reaction ..................................................................1111 B. Carbon–Nitrogen Bond–Forming Reaction ................................................................1112 VIII. N-Protected Enamines.........................................................................................................1112 IX. Hydroxylamines ..................................................................................................................1114 X. Carboxamides......................................................................................................................1116 XI. Hydroxamic Acids ...............................................................................................................1117 References .....................................................................................................................................1117

I. INTRODUCTION The electrochemical oxidation of nitrogen compounds is useful for the syntheses of complex organic molecules starting from easily available organic molecules. The oxidation reverses the polarity of functional groups, and therefore, molecules with nucleophilic nitrogens can be converted into electrophiles. Also, the oxidation selectively functionalizes starting amines and amino acids to afford building blocks for constructing a variety of more complex structures. The result has simplified the construction of synthetically useful nitrogen-containing compounds.

1103

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1104

Organic Electrochemistry

Many of the electrochemical oxidation reactions have excellent characteristics different from the reactions using traditional chemical oxidants. For this reason, the electrochemical oxidation of nitrogen-containing compounds is a valuable tool for any synthetic chemist to have at their disposal. This chapter majorly describes recent progress for electrochemical oxidations of aliphatic amines, N-protected aliphatic amines, N-protected amino alcohols, N-protected amino acids, hydroxylamines, carboxamides, and hydroxamic acids from a synthetic point of view. Accordingly, some books in these areas might help readers to comprehend important concepts and results sufficiently [1–3]. The electrochemical oxidation of easily oxidized nitrogen-containing compounds having lonepair electrons has been studied quite intensively. The selectivity of the reaction depends on the structures of such compounds. For example, electrochemical oxidation of aliphatic amines in many cases leads to a variety of products while the electrochemical oxidation of amides or carbamates results in the formation of only one major product [1,2].

II.

ALIPHATIC AMINES

A. PRIMARY AMINES The oxidation potentials of aliphatic amines depend on their structure. Primary amines are more difficult to oxidize than secondary or tertiary ones. The electrochemical oxidation of simple primary amines is usually quite complex and may lead to a variety of products. Among them, some methods for the transformation of primary amines into the corresponding nitriles were developed. Direct electrochemical oxidation using a nickel hydroxide electrode [4], and also indirect ones using 2,2,6,6-tetramethylpiperidinyl-1-oxyl (TEMPO) [5], or halogen ions (Scheme 28.1) [6] as mediators effectively transformed primary amines into the corresponding nitriles. In the latter reaction, anodically generated “Br+” promoted the oxidation in cooperation with cathodically generated MeO −.

B.

SECONDARY AND TERTIARY AMINES

Indirect electrochemical oxidation using halogen ions as mediator effectively transformed cyclic amino esters into relatively stable imines or enamines (Equation 28.1) [6]. –2e– CO2Me

N H

or N H

CO2Me or

N

NaBr/MeOH

CO2Me

N CO2Me H 67%

80%

–4e– NaBr/MeOH

RCH2NH2

MeO–

–2e–

Br+

RCN

Br–

Br–

–2e–

MeOH RCH

RCH2NHBr MeO–

4MeOH

SCHEME 28.1

+4e–

R = alkyl, aralkyl: 50–95%

MeOH

Br+

NH MeO–

MeO– RCH

NBr

MeOH

4MeO– + 2H2

Indirect electrochemical oxidation of primary amines to nitriles.

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(28.1)

1105

Aliphatic Nitrogen–Containing Compounds

Ph

–2e– NaOMe (6 equiv) LiClO4/MeOH

+ NaCN (6 equiv)

N Me

MeO N

MeO

Ph Me

NaBH4

+

N

H2

Ph

N Ph

CN

N H

Me

85%

85%

2) RX

CN Me

NH

MeO

1) LDA

Ph

N

MeO

MeO MeO

N

Ph

NC R Me

67% yield, 80% de RX = Mel 4-MeO-PhCH2Br 77% yield, 80% de 3,4-(MeO)2-PhCH2Br 75% yield, 80% de

R

R Me 76% yield, 98% de R = Me 85% yield, 80% de 4-MeO-PhCH2 3,4-(MeO)2-PhCH2 88% yield, 80% de

SCHEME 28.2

CN 6%

Me

MeO –2e– NaCN (2.5 eqiv) AcOH (0.5 equiv) MeO LiClO4/MeOH

MeO MeO

+

N Ph

Electrochemical cyanation of tert-aliphatic amines.

Electrochemical oxidation of sec- or tert-aliphatic amines in many cases led to a complex mixture or decomposed products while α-cyanation of cyclic amines proceeded selectively (Scheme 28.2) [7]. Electrochemical oxidation of tert-aminoalkylmalonates or tert-aminoalcohols smoothly proceeded to afford relatively stable cyclized products (Equations 28.2 and 28.3) [8]. CO2Me N

CO2Me

–2e– NaCN/MeOH

CO2Me N

CO2Me

(28.2)

73%

–2e– N HO

III.

N

Kl (0.25 equiv) NaOMe (1.25 equiv) MeOH

O

(28.3)

69%

N-PROTECTED AMINES

A. REGIOSELECTIVITY The electrochemical oxidation of amides, lactams, carbamates, and N-acylated amino acids is of considerable synthetic value because in contrast to aliphatic amines, the N,O-acetals formed by α-oxidation [9–11] are quite stable precursors of N-acyliminium ions and can be used effectively in amidoalkylation reactions (Scheme 28.3) [12,13]. In such amidoalkylations, the use of a chiral Lewis acid afforded optically active substituted products (Equation 28.4) [14]. n N PG

Chiral Lewis acid OMe

Nu–

n N

*

(28.4) Nu

PG Up to 97% ee

n = 0–2 PG: protecting group Nu: active methylene compounds, silyl enol ether, etc.

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1106

Organic Electrochemistry

R1 R2

OMe

–2e– MeOH

NCOR4

Lewis acid

R1

Nu

NCOR4

–

R3

R3 R1

NCOR4 R2

R2

R3

Nu R1

+

NCOR4 Nu: Ar, Het-Ar, alkyl, allyl, CN, PO(OR)2, etc. 2

R3

R

SCHEME 28.3

Electrochemical oxidation of N-protected amines and successive amidoalkylation. OH F3C N H

OH

OEt N

p-TsOH in benzene reflux

SiMe3

CF3

O

R

–2e– Graphite electrodes Et4NBF4, MeOH

n = 1–2 R: Alkyl, Ph, Allyl

SCHEME 28.5

CF3

+

N O CF3

N OH H

Electrochemical synthesis of (S)-α-allylprolinol.

( )n N CN

CF3CO2H in CH2Cl2

O

65%

10% HCl in MeOH

CF3

SCHEME 28.4

N

O

N

75%

MeO

–2e– MeOH

( )n N CN Major

OMe R

( )n Lewis acid Nu–

N CN

Nu R

Up to 100% regioselectivity

Electrochemical oxidation of N-cyano cyclic amines.

Also, excellent methods for oxidation and/or amidoalkylation of carbamates, such as the cation pool method, the cation flow method, recyclable solid supported bases, and parallel electrosynthesis (for these techniques, see Chapter 9), were developed [15,16]. As shown in Scheme 28.3, direct electrochemical oxidation occurred at the less substituted carbon, while some methods for electrochemical oxidation at the more substituted carbon were developed. Namely, electrochemical oxidation of a bicyclic amine prepared from (S)-prolinol and trifluoroacetaldehyde proceeded to afford an enantiomerically pure methoxylated compound in excellent regioselectivity. This product was easily transformed into (S)-α-allylprolinol (Scheme 28.4) [17]. Also, N-cyano-substituted cyclic amines were regioselectively methoxylated at the more substituted carbon by electrochemical oxidation (Scheme 28.5) [18]. Such opposite selectivity might be explained by the stability of the corresponding intermediary iminium ions.

B.

ELECTROCHEMICAL CYANATION

There are several publications available on electrochemical cyanation. However, most of the reactions reported suffer from low yield and long reaction time due to the use of a divided cell, several reaction steps, or specialized starting material [19]. Pilli and Santos published their work [20] on

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1107

Aliphatic Nitrogen–Containing Compounds ( )n

NC R

N

R

N

CN

PG

N

R

–2e–

–2e– Me3SiCN MeO2C

( )n

n = 1–4 PG = CO2Me, COPh, etc.

( )n

n = 1–2 PG = CN –2e–

MeO2C

Et4NBF4, MeSO3H CH2Cl2/MeCN

5 N

CN

N

CN

PG

92% yield cis/trans = 100:0

Tr

Tr Tr = triphenylmethyl

SCHEME 28.6 Regio- and/or diastereoselective cyanation.

electrochemical cyanation using two methods. In the case of the cation pool method [21] using a combination of TMSCN and TMSOTf, they achieved high yield and enantioselectivity; on the other hand, the use of the noncation pool electrochemical method using TMSCN gave very low yield and required low temperatures (−78°C). In addition, Tajima et al. have published their work on electrochemical cyanation based on the concept of site isolation [22]. The yields using this method were moderate to high, while the current efficiency was somewhat low. Recently, a highly efficient direct cyanation of N-protected cyclic amines by noncation pool electrochemical oxidation was reported [23]. This electrochemical cyanation of l-proline derivatives proceeded to afford 5-cis-substituted products in excellent diastereoselectivity (Scheme 28.6).

C.

ELECTROCHEMICAL α,β-FUNCTIONALIzATION

In relatively acidic reaction media, electrochemical α,β-functionalization of N-protected cyclic amines was achieved (Scheme 28.7) [24]. In these reactions, α-alkoxylated intermediates are somewhat unstable and easily convert into α,β-unsaturated cyclic amines, which might be successively oxidized. Such anodic diacetoxylation was applied to the kilogram-scale production of cis3-methylamino-4-methylpiperidines [24d].

D. PREPARATION OF IMIDES Indirect electrochemical oxidation mediated by N-hydroxyphthalimide was applicable to amides to afford imides (Scheme 28.8) [25]. In this oxidation, electrochemically generated phthalimide N-oxyl abstracts a hydrogen atom at the α-position of the starting amide to afford a radical intermediate, which was successively oxidized to imide. OAc

Cl n = 1–2 –4e–

( )n R

OR

N

n=2 –4e–

( )n R

NH4Cl/ROH

N

PG

AcOK/AcOH

( )n R

PG

PG –2e–

–2e–

ROH Cl–

( )n

Cl+ –2e–

R

N PG

SCHEME 28.7 Electrochemical α,β-functionalization.

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( )n

Acidic media OR

N

R

N PG

OAc

1108

Organic Electrochemistry R2 R1 O

–e– NHPI (0.25 equiv), pyridine (0.25 equiv)

N

R2 R1

0.85 V vs. SCE 0.1 M NaClO4 in MeCN

R3

O

N

O

R3

O

O N O

–e

–

N OH NHPI O

PINO O

R2 R1

N

O (MeCO)2NH: 60%

O : 60%

N O

R3 O

Me

N Me

O : 81%

SCHEME 28.8 Oxidation of amides to imides.

IV. N-PROTECTED α-AMINO ACIDS A.

CHEMOSELECTIVITY

Discrimination of similar functional groups in a molecule is applicable to a chemoselective reaction. Since there was a big difference between direct and indirect (chloride ion as mediator) electrochemical oxidation with respect to chemoselectivity, two regioisomeric unsaturated cyclic amines were obtained from an l-lysine derivative (Scheme 28.9) [26]. Indirect electrochemical oxidation of N-tosylamino compounds using halogen ions as a mediator in the presence of additional base promoted the migration of N-tosylamino groups via enamides to afford the corresponding aminoacetals (Scheme 28.10) [27].

B.

ENANTIOSELECTIVITY

A decarboxylative methoxylation of an N-acylated amino acid (Hofer–Moest reaction, non-Kolbe reaction, see Chapter 33) leads to N-acyl-iminium ion intermediates [28]. Although the transformation of an optically active α-amino acid into an active intermediate without any loss of optical purity is useful OMe CO2Me NH NH—CO Me 2 CO2Me

N

–2e– NaCl, MeOH

CO2Me NH NH—CO Me 2 CO2Me

Et4NOTs, MeOH

MeO

CO2Me NH NH—CO Me 2 CO2Me

N

CO2Me

CO2Me

CO2Me

CO2Me

SCHEME 28.9

–2e–

Chemoselective oxidation.

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1109

Aliphatic Nitrogen–Containing Compounds

Tos N H

HN Tos

–4e– KBr (0.5 equiv), KOH (0.5 equiv), MeOH

H N Tos

N

OMe

Tos

91%

OMe N Tos

H N

Tos

N

Tos

H N

H N

Tos

Ac

CO2Me

H N N H

Tos

Tos NH –4e– MeO KI (0.5 equiv), NaOMe (0.5 equiv), OMe MeOH

OMe

Br+ MeOH

Tos

92%

N H

H N Tos

Br

NH4Cl

H N

Tos

base

Br– –2e–

–2e–

Tos

H N

H N

Tos

Δ

Ac CO2Me

CO2Me

N Ac

82%

SCHEME 28.10 Oxidative migration of N-tosylamino groups. –2e–

CO2H

R1 *

Nu–

+

NCOR3

NCOR3

R2

Nu

R1

NCOR3 R2

R2 sp2 carbon intermediate

chiral

SCHEME 28.11

R1

racemic

Usual electrochemical decarboxylative substitution of N-acyl α-amino acids.

for the synthesis of optically active nitrogen-containing compounds, the intermediary iminium ion, which is a typical sp2 cation, might lose the original chirality to afford racemic products (Scheme 28.11). However, when N-o-phenylbenzoylated oxazoline and thiazoline derivatives were electrochemically oxidized, the memory of chirality via carbenium ion chemistry occurred to afford optically active products (83% and 91% enantiomeric excess [ee], respectively, in Equation 28.5) [29,30].

Me S Me N

Me O Me

N O

COOH or

O

Me Me COOH

Me O –2e– NaOMe MeOH –30°C

Me

S

Me N

OMe

Me

N

Me Me OMe

(28.5)

O

O

83% ee

91% ee

Scheme 28.12 shows a plausible stereochemical course for the memory of chirality. The initial step involves the oxidative decarboxylation of the amino acid to form the iminium ion, which can be attacked by nucleophiles (MeO −) from the syn or the anti side. The observed 85% ee could be attributed to the presence of the bulky o-phenyl group beneath the carboxylic group and the fixation of the conformation of the amino acid and the iminium ion intermediate at low temperature. The restricted rotation could favor the formation of a chiral iminium ion with the conformation of an o-phenyl group similar to that of the amino acid. The bulky o-phenyl group could preclude an effective approach from the anti side due to the steric repulsion between the o-phenyl group and MeO −, and hence, the nucleophilic attack was predominantly from the less hindered syn side resulting in the 4R-isomer.

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1110

Organic Electrochemistry OMe S

Me

MeO– CO2H S

Me N Me O

syn Me

–2e– NaOMe

N O Me

Pt cathode Pt anode MeOH, –30°C

+

n

te

Re

N Me O

n tio

S

4R-isomer 92.5% Inversion

Me

anti

S

N Me O

MeO–

OMe

: Steric repulsion 4S-isomer 7.5% Chiral iminium ion

SCHEME 28.12 Plausible stereochemical course for the memory of chirality.

V. N-PROTECTED β-AMINO ALCOHOLS Electrochemical oxidation of N-acyl-β-amino alcohols smoothly cleaves the carbon–carbon bond to afford N,O-acetals [31]. A memory of chirality was observed in the electrochemical substitution of optically active β-amino alcohol derivatives (Equation 28.6) [32]. α N O

β CAr2 OH

R N

–2e– Pt electrodes

OMe

Ar

O

NaOMe (1.2 equiv) in MeOH, at –30°C

Ph

Ar

(28.6)

+ O

Ph

Up to 73% ee

On the other hand, indirect electrochemical oxidation in the presence of a chiral copper catalyst transformed racemic N-protected aminoalcohols into optically active amino esters by kinetic resolution (Scheme 28.13) [33]. Similar kinetic resolution of racemic N-protected aminoalcohols proceeded to afford optically active amino esters. In this reaction, chelation of the amino alcohol or the amino aldehyde with the Lewis acid activates their hydroxyl or formyl group to form alkoxide ion, which is easily oxidizable compared with the original amino alcohol or aldehyde (Scheme 28.14) [34]. –4e– Pt electrodes, rt Et4NBr (1.0 equiv)

OH N Bz

N Bz

MeOH Cu(OTf )2 (0.1 equiv) (R,R)-Ph-BOX (0.1 equiv)

racemic

OH

COOMe

N Bz

+

(R)-ester

(S)-alcohol

27% yield 70% ee

50% yield 15% ee

O

O N

N Bz

CHO

rac-16–19

SCHEME 28.13

–2e– Pt electrodes, rt Et4NBr (1.0 equiv) MeOH Cu(OTf )2 (0.1 equiv) (R,R)-Ph-BOX (0.1 equiv)

N Bz

COOMe +

N Bz

CHO

(R)-ester

(S)-alcohol

43% yield 86% ee

34% yield 27% ee

Enantioselective oxidation of amino alcohol derivatives.

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N

Ph Ph L* (R,R)-Ph-BOX

1111

Aliphatic Nitrogen–Containing Compounds R4 R3 N Bz

R4

Cu2+

N O H Bz Cu2+ L*: (R,R)-Ph-BOX

MeO

R3

Br

L*

SCHEME 28.14

R3

N * CHO Bz

L*

L*: (R,R)-Ph-BOX

MeO–

L*

Cathode

N* Bz

COOMe

MeOH + Br–

R4 OMe MeO–

OMe R

N O– Bz Cu2+

2MeO– + H2

R4 R3

L*

R4

2e–

CHO Cu2+

L*

–

2MeOH

N Bz

N O Bz Cu2+

R3

O Br N Bz Cu2+

N O– Bz Cu2+

H

MeO–

R3 +

VI.

R3 MeOH + Br

R4

R4

MeOH

R4

R4

L*

–

R3

Anode 2e–

R3

L*

R4

Br+

Br–

OH

Br+

3

N O Br Bz Cu2+ L*

Reaction mechanism for enantioselective oxidation of amino alcohols or aldehydes.

N-PROTECTED α-ALLyL OR bENZyL AMINES

Electrochemical oxidation of N-acyl-α-allyl or benzyl amines smoothly cleaves the carbon–carbon bond to afford N,O-acetals. The allyl groups work as chiral auxiliary to afford optically active quaternary cyclic amino acids (Scheme 28.15) [35]. Similarly, electrochemical oxidation of α-silyl-or α-thio-substituted carbamates or azetidin-2-ones smoothly proceed to afford the corresponding N,O-acetals [36,37].

VII. INTRAMOLECULAR ANODIC COUPLINg OF N-PROTECTED AMINES A.

CARBON–CARBON BOND –FORMING REACTION

An important intermediate for the preparation of carbapenem antibiotics was synthesized by an indirect electrochemical intramolecular carbon–carbon bond–forming reaction and direct electrochemical decarboxylative methoxylation (Scheme 28.16) [38]. In this cyclization, the (R)-phenylethyl group works as a good chiral auxiliary.

–2e– N Cbz

N CO2Me Cbz

or

N Cbz

Ph

THF –78°C to rt 81%

or

N OMe Cbz

76%

–2e– Et4NBF4 MeO MeOH 0°C 92%

NaHMDS (1.2 equiv) MeI (3.0 equiv)

N OMe Cbz

Et4NBF4 in MeOH–MeCN

74%

Allyl–TMS (1.2 equiv) BF3· OEt2 (1.2 equiv) N CO2Me Cbz

Me N CO2Me Cbz

CH2Cl2 –78°C to rt 88%

–2e– Et4NBF4 MeOH–MeCN –10°C 79%

MeO

N CO2Me Cbz cis isomer

Me N CO2Me Cbz

Et3SiH (1.5 equiv) MeSO3H (1.2 equiv) CH2Cl2 –78°C to rt 90%

Me N CO2Me Cbz >99% ee

SCHEME 28.15 Electrochemical deallylation or debenzylation of α-allyl- or benzyl-cyclic amine derivatives.

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1112

Organic Electrochemistry O

O

CO2-t-Bu CO2-t-Bu

Me O Me

–2e– NaI (0.5 equiv)

N

85°C in MeCN

R

94% yield 80% de

OAc R H Me S

CO2H

N O Me R

N O Me R

Me

O Me

R

OH H S N O Me R

83% yield 84% de

2) Chromatography

CO2-t-Bu CO2-t-Bu

OSiMe2-t-Bu R H OMe Me R S NH O

OMe R N R

63% yield

4R:4S = 100:0

73% yield

–2e– KI MeOH

I CO Me 2 n

CO2Me NH

EGB

n

CO2H

CO2Me CO2Me

O n

n

N Tos

N

N Tos

Tos

n=1–3

SCHEME 28.17

CO2-t-Bu

OH R H Me S

1) –2e–, NaOMe MeOH/MeCN (1/4)

NaBH4 in THF

Electrochemical carbon–carbon bond–forming reaction between active methylene groups.

CO2Me n CO2Me NH Tos

CO2-t-Bu

S

75% yield

SCHEME 28.16

H

Me

Tos

Up to 100% yield

Coupling of nitrogen and active methylene.

OMe EDG

( )n

–e–, RVC anode LiOMe, Et4NOTs

NH Tos

MeOH-THF n = 2–4

SCHEME 28.18

B.

( )n

+ N–

EDG

Tos

EDG

( )n

EDG

( )n

N

N

Tos

Tos Up to 91% yield

Coupling of nitrogen and alkene.

CARBON–NITROGEN BOND –FORMING REACTION

The electrochemical intramolecular carbon–nitrogen bond–forming reaction of N-tosyl-aminoalkylmalonates smoothly proceeded to afford nitrogen heterocycles (Scheme 28.17) [39]. Electrogenerated radical cations from electron-rich alkenes were intramolecularly trapped with nitrogen to afford nitrogen heterocycles (Scheme 28.18) [40].

VIII.

N-PROTECTED ENAMINES

Since N-protected enamines are representative electron-rich olefins, they are relatively easy to oxidize. Direct electrochemical oxidation of a 6-acetoxymethyl-2,3-didehydropiperidine derivative afforded the 3,6-trans isomer, whereas indirect oxidation gave the 3,6-cis isomer in high diastereoselectivity. On the other hand, the indirect method using I− as a mediator proceeded via inversion of the stereochemistry (Scheme 28.19) [41].

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1113

Aliphatic Nitrogen–Containing Compounds Direct electrochemical oxidation AcO OAc

N CO2Bn

AcOK AcOH

AcO

AcO

Et3SiH

–2e– OAc

N

3

MeSO3H CH2Cl2

CO2Bn

6

OAc

N

CO2Bn 91% yield cis/trans = 32/68

Indirect electrochemical oxidation HO OAc

N CO2Bn

SCHEME 28.19

OAc

N

HO

NaI (0.1 equiv) H2O–acetone (1:9)

HO

Et3SiH

–2e–

3

MeSO3H CH2Cl2

CO2Bn

6

OAc

N

CO2Bn 72% yield cis/trans = >99/99% ee

quant.

SCHEME 28.26 Preparation of a chiral N-oxyl.

X. CARbOXAMIDES Electrolysis of carboxamides in MeOH containing bromide ions efficiently led to products of the Hofmann rearrangement. This reaction, named the electrochemically induced (E-I) Hofmann rearrangement, is achieved without any bromine and base under mild and neutral reaction conditions. Thus, for example, an epoxy functional group in the alcohol part remained intact during the conversion of carboxamides to carbamates (Scheme 28.27) [52]. However, the conditions for the E-I Hofmann rearrangement are not yet suitable for substrates that are unstable under weakly basic conditions since EGBs may be present in the vicinity of the cathode. For example, the E-I Hofmann rearrangement in MeOH containing bromide ions yielded the desired Hofmann rearrangement product in only a low yield (28%) and a by-product that might be generated by the base-catalyzed cyclization of the starting amide. On the other hand, the transformation of carboxamides to carbamates was successfully achieved without any formation of imides and no loss of the optical purity by electrolysis using the CF3CH2OH/MeCN solvent system in which CF3CH2OH might play an important role to control the basicity caused by the EGB (Scheme 28.28) [53].

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1117

Aliphatic Nitrogen–Containing Compounds

O

O NH2

+

–2e– Et4NBr (0.5 equiv)

HO

MeCN 53%

EGB

–2e–

Br+

O

H N

O O

Br–

O HO

EGBH+

EGBH+

EGB

O

O – NBr

NHBr

NCO

SCHEME 28.27 The E-I Hofmann rearrangement. NHBoc O

CO Me NH2 2

CO Me NH2 2

NHBoc

NHBoc NH CO2Me CO2Me 28% (>99.9% ee)

MeOH–MeCN at 60°C

NHBoc O

–2e– Et4NBr (0.5 equiv)

–2e– Et4NBr (0.5 equiv)

+ O

N H

14% (25% ee)

NHBoc NH CO2Me CO2CH2CF3

CF3CH2OH–MeCN at 60°C

O

NHBoc +

80% (>99.9% ee)

O

O N H 0%

SCHEME 28.28 The E-I Hofmann rearrangement under neutral conditions. +R2NH2 –2e– R1CO

N OH R

R1CONHR2

R1CO+ –O –RN –

+R3OH

R1COOR3

SCHEME 28.29 Electrochemical oxidation of hydroxamic acids.

XI. HyDROXAMIC ACIDS Electrochemical oxidation of hydroxamic acids in the presence of amines or alcohols afforded the corresponding amides or esters (Scheme 28.29) or carboxylic acids by reaction with water [54].

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Organic Electrochemistry

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Scand. 1992, 46, 194–199. (d) McClure, K. F.; Renold, P.; Kemp, D. S. J. Org. Chem. 1995, 60, 454–457. (e) Danielmeier, K.; Schierle, K.; Steckhan, E. Angew. Chem. Int. Ed. 1996, 35, 2247–2248. (f) Dhimane, H.; Vanucci-Bacqué, C.; Hamon, L.; Lhmmet, G. Eur. J. Org. Chem. 1998, 1955–1963. (g) Grossmith, C. E.; Senia, F.; Wagner, J. Synlett 1999, 10, 1660–1662. (h) D’Oca, M. G. M.; Pilli, R. A.; Vencato, I. Tetrahedron Lett. 2000, 41, 9709–9712. (i) Kinderman, S. S.; van Maarseveen, J. H.; Schoemaker, H. E.; Hiemstra, H.; Rutjes, F. P. J. T. Synthesis 2004, 36, 1413–1418. (j) Bartels, M.; Zapico, J.; Gallagher, T. Synlett 2004, 15, 2636–2638. (k) Sierecki, E.; Turcaud, S.; Martens, T. Royer, J. Synthesis 2006, 38, 3199–3208. (l) Bodmann, K.; Bug, T.; Steinbeisser, S.; Kreuder, R.; Reiser, O. Tetrahedron Lett. 2006, 47, 2061–2064. 12. (a) Shono, T.; Matsumura, Y.; Tsubata, K. J. Am. Chem. Soc. 1981, 103, 1172–1176. (b) Shono, T.; Matsumura, Y.; Tsubata, K. Org. Synth. 1985, 63, 206–213. (c) Shono, T.; Matsumura, Y.; Uchida, K.; Kobayashi, H. J. Org. Chem. 1985, 50, 3243–3245. (d) Matsumura, Y.; Kanda, Y.; Shirai, K.; Onomura, O.; Maki, T. Org. Lett. 1999, 1, 175–178. (e) Matsumura, Y.; Onomura, O.; Suzuki, H.; Furukubo, S.; Maki, T.; Li, C.-J. Tetrahedron Lett. 2003, 44, 5519–5522. (f) Matsumura, Y.; Ikeda, T.; Onomura, O. Heterocycles 2006, 67, 113–117. (g) Onomura, O.; Kirira, P. G.; Tanaka, T.; Tsukada, S.; Matsumura, Y.; Demizu, Y. Tetrahedron 2008, 64, 7498–7503. (h) Kamogawa, S.; Ikeda, T.; Matsumura, Y.; Kuriyama, M.; Onomura, O. Heterocycles 2010, 82, 325–332. (i) Hirata, S.; Kuriyama, M.; Onomura, O. Tetrahedron 2011, 67, 9411– 9416. (j) Mizuta, S.; Onomura, O. RSC Adv. 2012, 2, 4850–4853. 13. (a) Speckamp, W. N.; Moolenaar, M. J. Tetrahedron 2000, 56, 3817–3856. (b) Yazici, A.; Pyne, S. G. Synthesis 2009, 41, 339–368. (c) Yazici, A.; Pyne, S. G. Synthesis 2009, 513–541. (d) de Koning, H.; Speckamp, W. N. In: Helmchen, G.; Hoffmann, R. W.; Mulzer, J.; Schaumann, E., eds.; Houben-Weyl, Stereoselective Synthesis, Vol. E21, Georg Thieme, Stuttgart, Germany, 1995, pp.  1953–2009. (e) Hiemstra, H.; Speckamp, W. N. In: Trost, B. M.; Fleming, I.; Heathcock, C. H., eds.; Comprehensive Organic Synthesis, Vol. 2, Pergamon, Oxford, U.K., 1991, pp. 1047–1082. (f) Volkmann, R. A. In: Trost, B. M.; Fleming, I.; Schreiber, S. L., eds.; Comprehensive Organic Synthesis, Vol. 1, Pergamon, Oxford, U.K., 1991, pp. 355–396. 14. (a) Onomura, O.; Kanda, Y.; Nakamura, Y.; Maki, T.; Matsumura, Y. Tetrahedron Lett. 2002, 43, 3229– 3231. (b) Onomura, O.; Kanda, Y.; Imai, M.; Matsumura, Y. Electrochim. Acta 2005, 50, 4926–4935. (c) Onomura, O.; Ikeda, T.; Matsumura, Y. Heterocycles 2005, 66, 81–86. (d) Minato, D.; Imai, M.; Kanda, Y.; Onomura, O.; Matsumura, Y. Tetrahedron Lett. 2006, 47, 5485–5488. (e) Matsumura, Y.; Minato, D.; Onomura, O. J. Organomet. Chem. 2007, 692, 654–663. 15. (a) Yoshida, J.; Suga, S.; Suzuki, S.; Kinomura, N.; Yamamoto, A.; Fujiwara, K. J. Am. Chem. Soc. 1999, 121, 9546–9549. (b) Suga, S.; Okajima, M.; Fujiwara, K.; Yoshida, J. J. Am. Chem. Soc. 2001, 123, 7941–7942. (c) Horii, D.; Fuchigami, T.; Atobe, M. J. Am. Chem. Soc. 2007, 129, 11692–11693. (d) Yoshida, J.; Kataoka, K. Horcajada, R.; Nagaki, I. Chem. Rev. 2008, 108, 2265–2299. 16. (a) Tajima, T.; Fuchigami, T. Chem. Eur. J. 2005, 11, 6192–6196. (b) Siu, T.; Li, W.; Yudin, A. K. Comb. Chem. 2000, 2, 545–549. 17. (a) Onomura, O.; Ishida, Y.; Maki, T.; Minato, D.; Demizu, Y.; Matsumura, Y. Electrochemistry 2006, 74, 645–648. (b) Dhimane, H.; Vanucci, C.; Lhommet, G. Tetrahedron Lett. 1997, 38, 1415–1418. 18. Libendi, S. S.; Demizu, Y.; Matsumura, Y.; Onomura, O. Tetrahedron 2008, 64, 3935–3942. 19. (a) Chiba, T.; Takata, Y. J. Org. Chem. 1977, 42, 2973–2977. (b) Konno, A.; Fuchigami, T.; Fujita, Y.; Nonaka, T. J. Org. Chem. 1990, 55, 1952–1954. (c) Gall, E. L.; Hurvois, J. -P.; Sinbanbhit, S. Eur. J. Org. Chem. 1999, 2645–2653. (d) Girard, N.; Hurvois, J.-P. Tetrahedron Lett. 2007, 48, 4097–4099. 20. Shankaraiah, N.; Pilli, R. A.; Santos, L. S. Tetrahedron Lett. 2008, 49, 5098–5100. 21. Yoshida, J.; Suga, S. Chem. Eur. J. 2002, 8, 2650–2658. 22. Tajima, T.; Nakajima, A. J. Am. Chem. Soc. 2008, 130, 10496–10497. 23. Libendi, S. S.; Demizu, Y.; Onomura, O. Org. Biomol. Chem. 2009, 7, 351–356.

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Aliphatic Nitrogen–Containing Compounds

1119

24. (a) Shono, T.; Matsumura, Y.; Onomura, O.; Kanazawa, T.; Habuka, M. Chem. Lett. 1984, 13, 1101–1104. (b) Shono, T.; Matsumura, Y.; Onomura, O.; Ogaki, M.; Kanazawa, T. J. Org. Chem. 1987, 52, 536–541. (c) Shono, T.; Matsumura, Y.; Onomura, O.; Sato, M. J. Org. Chem. 1988, 53, 4118–4121. (d) Cai, W.; Colony, J. L.; Frost, H.; Hudspeth, J. P.; Kendall, P. M.; Krishnan, A. M.; Makowski, T.; Mazur, D. J.; Phillips, J.; Ripan, D. H. B.; Ruggeri, S. G.; Stearns, J. F.; White, T. D. Org. Process Res. Dev. 2005, 9, 51–56. 25. Masui, M.; Hara, S.; Ozaki, S. Chem. Pharm. Bull. 1986, 34, 975–979. 26. (a) Shono, T.; Matsumura, Y.; Inoue, K. J. Org. Chem. 1983, 48, 1388–1389. (b) Shono, T.; Matsumura, Y.; Inoue, K. J. Chem. Soc. Chem. Commun. 1983, 19, 1169–1171. 27. (a) Shono, T.; Matsumura, Y.; Katoh, S.; Inoue, K.; Matsumoto, Y. Tetrahedron Lett. 1986, 27, 6083–6086. (b) Shono, T.; Matsumura, Y.; Katoh, S.; Takeuchi, K.; Sasaki, K.; Kamada, K.; Shimizu, R. J. Am. Chem. Soc. 1990, 112, 2368–2372. 28. (a) Iwasaki, T.; Horikawa, H.; Matsumoto, K.; Miyoshi, M. J. Org. Chem. 1979, 44, 1552–1554. (b) Shono, T.; Matsumura, Y.; Tsubata, K.; Uchida, K. J. Org. Chem. 1986, 51, 2590–2592. (c) Zietlow, A.; Steckhan, E. J. Org. Chem. 1994, 59, 5658–5661. 29. (a) Wanyoike, G. N.; Onomura, O.; Maki, T.; Matsumura, Y. Org. Lett. 2002, 4, 1875–1877. (b) Wanyoike, G. N.; Matsumura, Y.; Kuriyama, M.; Onomura, O. Heterocycles 2010, 80, 1177–1185. 30. (a) Matsumura, Y.; Shirakawa, Y.; Satoh, Y.; Umino, M.; Tanaka, T.; Maki, T.; Onomura, O. Org. Lett. 2000, 2, 1689–1691. (b) Matsumura, Y.; Tanaka, T.; Wanyoike, G. N.; Maki, T.; Onomura, O. J. Electroanal. Chem. 2001, 507, 71–74. (c) Matsumura, Y.; Wanyoike, G. N.; Onomura, O.; Maki, T. Electrochim. Acta 2003, 48, 2957–2966. 31. Shono, T.; Matsumura, Y.; Tsubata, K.; Sugihara, Y. Nippon Kagaku Kaishi 1984, 1782–1787. 32. Wanyoike, G. N.; Matsumura, Y.; Onomura, O. Heterocycles 2009, 79, 339–345. 33. Minato, D.; Arimoto, H.; Nagasue, Y.; Demizu, Y.; Onomura, O. Tetrahedron 2008, 64, 6675–6683. 34. (a) Onomura, O.; Arimoto, H.; Matsumura, Y.; Demizu, Y. Tetrahedron Lett. 2007, 48, 8668–8672. (b) Minato, D.; Nagasue, Y.; Demizu, Y.; Onomura, O. Angew. Chem. Int. Ed. 2008, 47, 9458–9461. (c) Maki, T.; Iikawa, S.; Mogami, G.; Harasawa, H.; Matsumura, Y.; Onomura, O. Chem. Eur. J. 2009, 15, 5364–5370. 35. Kirira, P. G.; Kuriyama, M.; Onomura, O. Chem. Eur. J. 2010, 16, 3970–3982. 36. Sugawara, M.; Mori, K.; Yoshida, J. Electrochim. Acta 1997, 42, 1995–2003. 37. Suda, K.; Hotoda, K.; Iemuro, F.; Takanami, T. J. Chem. Soc. Perkin Trans. 1 1993, 1553–1555. 38. Minato, D.; Mizuta, S.; Kuriyama, M.; Matsumura, Y.; Onomura, O. Tetrahedron 2009, 65, 9742–9748. 39. Shono, T.; Matsumura, Y.; Katoh, S.; Ohshita, J. Chem. Lett. 1988, 17, 1065–1068. 40. Xu, H.-C.; Moeller, K. D. J. Am. Chem. Soc. 2010, 132, 2839–2844. 41. Libendi, S. S.; Ogino, T.; Onomura, O.; Matsumura, Y. J. Electrochem. Soc. 2007, 154, E31–E35. 42. (a) Furukubo, S.; Moriyama, N.; Onomura, O.; Matsumura, Y. Tetrahedron Lett. 2004, 45, 8177–8181. (b) Moriyama, N.; Matsumura, Y.; Kuriyama, M.; Onomura, O. Tetrahedron Asymmetry 2009, 20, 2677–2687. 43. Sayo, H.; Ozaki, S.; Masui, M. Chem. Pharm. Bull. 1973, 21, 1988–1995. 44. Sayo, H.; Ozaki, S.; Masui, M. Chem. Pharm. Bull. 1975, 23, 1702–1707. 45. Karady, S.; Corley, E. G.; Abramson, N. L.; Amato, J. S.; Weinstock, L. M. Tetrahedron 1991, 47, 757–766. 46. Shono, T.; Matsumura, Y.; Inoue, K. J. Org. Chem. 1986, 51, 549–551. 47. Semmelhack, M. F.; Chou, C. S.; Cortes, D. A. J. Am. Chem. Soc. 1983, 105, 4492–4494. 48. Kashiwagi, Y.; Kurashima, F.; Chiba, S.; Anzai, J.; Osa, T.; Bobbitt, T. M. Chem. Commun. 2003, 39, 114–115. 49. Tanaka, H.; Kawakami, Y.; Goto, K.; Kuroboshi, M. Tetrahedron Lett. 2001, 42, 445–448. 50. Shiigi, H.; Mori, H.; Tanaka, T.; Demizu, Y.; Onomura, O. Tetrahedron Lett. 2008, 49, 5247–5251. 51. Demizu, Y.; Shiigi, H.; Mori, H.; Matsumoto K.; Onomura, O. Tetrahedron: Asymmetry 2008, 19, 2659–2665. 52. (a) Matsumura, Y.; Maki, T.; Satoh, Y. Tetrahedron Lett. 1997, 38, 8879–8882. (b) Matsumura, Y.; Satoh, Y.; Maki, T.; Onomura, O. Electrochim. Acta 2000, 45, 3011–3020. 53. Matsumura, Y.; Satoh, Y.; Shirai, K.; Onomura, O.; Maki, T. J. Chem. Soc. Perkin Trans. 1 1999, 2057–2060. 54. Masui, M.; Ueshima, T.; Yamazaki, T.; Ozaki, S. Chem. Pharm. Bull. 1983, 31, 2130–2133.

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29

Aromatic Nitrogen–Containing Compounds Jan S. Jaworski

CONTENTS I.

Aniline and Its Derivatives ....................................................................................................1121 A. Unsubstituted Aniline ....................................................................................................1121 B. Nitrogen-Substituted Anilines ...................................................................................... 1124 C. Ring-Substituted Anilines............................................................................................. 1127 II. Other Aromatic Amines ....................................................................................................... 1136 III. Azo and Azoxy Compounds ................................................................................................. 1137 IV. Aryl Diazonium Salts ............................................................................................................1142 V. Aromatic Triazenes ................................................................................................................1143 Acknowledgment ..........................................................................................................................1145 References .................................................................................................................................... 1146 Of all aromatic amines, aniline and its derivatives have been most extensively studied electrochemically for many years. The important aim of all those studies was to explain the mechanism of formation of a number of products of this complex oxidation process performed under different electrolysis conditions and using reactants with various substituents. A lot of investigators have focused their interest on understanding the early stages of electrode reactions that yield polyanilines giving rise to many important applications. In view of these, the main part of this chapter is devoted to fundamental results concerning the mechanisms of anodic processes of unsubstituted aniline as well as N-substituted and C-substituted anilines. For the last two groups, monosubstituted anilines in para-, meta-, and ortho-positions are considered first, followed by the discussion of di- and tri-substituted anilines. A brief review of the electrochemical reactions of other aromatic amines including polyaminobenzenes follows, and the electrode behavior of azo and azoxy compounds, diazonium salts, as well as aromatic triazenes is reported at the end. In this chapter, only the main routes of the electrode processes and the most important products are considered on the basis of some examples, and only references to the selected examples are cited. A more comprehensive discussion and references to investigations of a number of particular compounds can be found in recent monographic chapters dealing with the electrochemistry of anilines [1] and azo compounds [2]. In earlier monographs, the half-wave and cyclic voltammetric (CV) peak potentials as well as products of electrode processes performed under different conditions are collected in tables for many aniline, diphenylamine, triphenylamine, and phenylenediamine derivatives [3], as well as for azo, diazo, and azoxy compounds [4].

I.

ANILINE AND ITS DERIVATIVES

A. UNSUBSTITUTED ANILINE The first electron transfer from aniline (I) produces the radical cation C6H5NH2+ • (II) that can be further oxidized to dication C6H5NH22+ (III). As usual, cations II and III are electrophilic and more acidic than a parent molecule and thus, they have a tendency to deprotonation yielding the aniline 1121

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1122

Organic Electrochemistry

C6H5NH2 (I)

–e E1

C6H5NH2+ (II)

H+

–e E2

–e (IV)

SCHEME 29.1

(III)

H+

pK1

C6H5NH

C6H5NH22+

E3

pK2

C6H5NH+ (V)

Equilibria of electron and proton transfers during the oxidation of aniline I.

neutral radical (IV) and the nitrenium cation (V), respectively. The square diagram involving all these electron and proton transfers is shown in Scheme 29.1 [5]. It explains why the whole oxidation process and the properties of polyaniline formed are strongly dependent on the medium nature, first of all on the acidity and the use of aqueous or nonaqueous solvents. No direct experimental evidence could be found for the formation of unstable radical cation II under acidic conditions, for example, in 1 M H2SO4 using fast CV measurements [6] or in a eutectic mixture NH4F + 2.3 HF using in situ ESR techniques [5]. However, in acetonitrile solution for low aniline concentrations (0.05  mM) using a platinum disk electrode of 25 μm diameter and a scan rate of 200 V s−1 the CV reduction peak of the radical cation II was observed for the first time by Yang and Bard [6]. Moreover, the high value (108 M−1 s−1) of the estimated rate constant for the further coupling of II explains unsuccessful attempts at ESR measurements under acidic conditions. Note however that radical cations and nitrenium ions of some substituted anilines are more stable as will be discussed at the end of Section I.C. The formation of the nitrenium cation (V) during the electrode oxidation of aniline in NH4F + 2.3 HF mixture was demonstrated as the only possible intermediate responsible for the formation of a polymer containing the phenazine ring, which gives a characteristic peak in CV curves of polyaniline [7]. However, Yang and Bard explained later [6] that the aforementioned peak corresponds to different redox couples depending on the aniline concentration and the experimental time scale. Nevertheless, the band observed at 420 nm in spectroelectrochemical measurements was attributed by Geniès et  al. to V [7]. Additional support of the square diagram shown in Scheme 29.1 was obtained using a rotating ring-disk platinum electrode where intermediates were collected at a ring [8]. The oxidation of 0.2 M aniline in an aqueous solution containing 0.5 M KCl showed two distinct peaks at 1.0 and 1.1 V versus the Ag/AgCl/saturated KCl reference electrode that correspond to potentials E1 and E3 in Scheme 29.1. However, only one intermediate II was formed in solutions of high acidity when the radical cation II does not deprotonate. It should be remembered that the molecules of all intermediates shown in Scheme 29.1 are superpositions of resonance structures and the individual structure with charge and spin densities at different nitrogen or carbon atoms can be considered in further coupling reactions. In particular, primary products of the decay of radical cations II in acidic aqueous media, namely benzidine (VI) and p-aminodiphenylamine (VII), are formed by tail-to-tail and head-to-tail coupling, respectively, as is shown in Scheme 29.2. These two major products found in aqueous H2SO4 solutions are formed by the radical cation–radical cation coupling in the DIM1 mechanism (Scheme 29.2) and not by the dimerization of the radical cation II with the neutral parent molecule I (the DIM2 mechanism). This conclusion was proven by digital simulations of CV curves obtained at a glassy carbon electrode [6]. However, the head-to-head coupling of radical cations II (Scheme 29.2) to form hydrazobenzene (VIII) can also be observed in acidic [9] and alkaline [10] solutions. In general, products of all reactions shown in Scheme 29.2, that is, tail-to-tail, tail-to-head, and head-to-head coupling, can exist at electrodes forming different mixtures of higher oligomers. This is very important for the properties of final polymers and their practical applications.

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1123

Aromatic Nitrogen–Containing Compounds

–

H N+ H –

H –2H+

H

–2e H2N

NH2

–

H

+

–

NH

(VI)

–

N

HN

–2H+

H

(XV)

+N

–

H

–2e

H –2H+

(II)

N H

(VII)

–

–2e N N

–2H+

– –

–

(II)

(VIII)

(II)

SCHEME 29.2

(XVII)

N N H H

–

–

H H N+ + + N H H

NH2+

N

–2H+

–

H 2N

(XVI)

Coupling reactions of aniline radical cations II.

The distribution of primary dimers VI, VII, and VIII strongly depends on the detailed conditions of electrolysis. Some typical examples are mentioned further on. The oxidation of aniline in 1 M H2SO4 at a platinum electrode yields VI and VII. However, VII, which is the dominant intermediate with the fastest reaction kinetics, is protonated in acidic media forming IX (Scheme 29.3), which can be further oxidized to N-phenylquinonediimine (X) in solutions having the pH of 1.2–4.8 [11]. Moreover, IX and X can comproportionate forming XI; the corresponding equilibrium constant in DMSO containing 0.1–0.5 M sulfuric acid was found to be Kcom = 0.027 [12]. Moreover, the formation of the radical cation XI in the reaction of II with I in the aqueous solution of 0.5 M H2SO4 was observed using in  situ time-resolved FT-IR spectroscopy [13]. It should also be noted that X undergoes a slow oxidative decomposition to N-phenylquinoneimine (XII) and next to the anilinium cation (XIII) and p-benzoquinone (XIV). Products XII and XIV may be reversibly reduced as is shown in Scheme 29.3 [6].

NH+3

N H

–2e

(IX)

+IX

NH+2

N

–2H+

NH+2

N H

2 Kcom

(X)

(XI)

H3O+ 2e

O

N

N H

2H+

(XII) H3O+ NH3+

O

OH 2e

+

(XIII)

SCHEME 29.3

Reaction pathways for cation IX.

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2H+ O (XIV)

OH

OH

1124

Organic Electrochemistry

In a similar manner, benzidine (VI) formed in the tail-to-tail coupling can undergo reversible oxidation [6] with the exchange of two electrons and two protons yielding (1,1′-biphenyl)-4,4′diimine (XV), as is shown in Scheme 29.2. It can be added that further oxidation of the dimers to radical cations and their recombination with monomer radical cations produce trimers that can be detected in solutions. However, higher oligomers are not formed in solution but on the electrode [14] and the electrode filming by oxidation products makes their identification rather difficult. On the other hand, the head-to-head coupling of the radical cations II (besides two other coupling modes discussed earlier) was observed in 0.1 M H2SO4 during the oxidation of aniline at graphite electrodes [9]. The hydrazobenzene (VIII) formed in the last process is next oxidized to azobenzene (XVI) as is shown in Scheme 29.2. However, VIII was not detected in 2.2 M sulfuric acid. In more recent papers using the electrochemical thermospray mass spectrometry (MS), the only deuterated dimer formed during the oxidation at a platinum foil of 0.01 M d5 -aniline in 0.1 M sulfuric acid was identified as IX [14] and similarly VII was identified during the oxidation of 10 −3 M aniline in aqueous solutions buffered at pH = 5 [15]. Moreover, interesting results were reported using a thin-layer electrochemical flow cell (with a glassy carbon electrode) coupled online with the electrospray MS in buffered aqueous solutions and in 1/1 (v/v) aqueous methanol mixtures. All solutions contained ammonium acetate and either acetic acid or ammonium hydroxide added to obtain a proper pH. In methanol mixtures, the following soluble dimeric products were found: VI, VII, and XVII at pH = 4, but VII, VIII, XVI, and XVII at pH = 9 [16]. Taking into account different competitions of electron transfer steps and homogeneous coupling reactions during the electrode oxidation of aniline, as revealed in Schemes 29.1 and 29.2, a complex pattern of multiple redox peaks can be observed in CV curves recorded in various conditions. It is worth noting and highly recommended that their comparison and interpretation should be made with great caution.

B.

NITROGEN-SUBSTITUTED ANILINES

In general, the nitrogen substitution of aniline results in the formation of more stable radical cations as intermediates of the oxidation than in the case of the unsubstituted reactant. Therefore, the main route of the electrode reaction is changed. The anodic oxidation of N-alkyl anilines in acidic aqueous media produces radical cations with the positive charge delocalized at the ring due to the inductive effect of the electron-releasing substituent in such a way that the tail-to-tail coupling is favored. The increase of steric hindrance by the nitrogen substitution also favors the –C–C– coupling. Thus, benzidine derivatives are the main products of dimerization, as will be shown further on. Protonated N,N′-dimethylbenzidine N(Me)H–C6H4 –C6H4 –N+(Me)H2 (XVIII) and protonated N,N′-dimethyl-p-aminodiphenylamine C6H4 –N(Me)–C6H4 –N+(Me)H2 (XIX) in the ratio of 7:3 are the early products of the oxidation of N-methylaniline in 0.1 M H2SO4. These products were identified by the electrochemical thermospray MS using the reactant deuterated at the N–H group [14]. However, the aforementioned authors also found that dimers undergo dealkylation through the formation of the –N+(H)=CH2 group by the loss of two electrons and a proton from the –N(H)Me group. Further hydrolysis with the oxidation yields the radical cation possessing the –N+ •H2 group. As a result, the dealkylated protonated dimers NH2–C6H4 –C6H4 –N+(Me)H2 and C6H4 –NH–C6H4 – N+(Me)H2 were also indentified. Moreover, trimers formed from XVIII by C–C–N–C coupling and two trimers from XIX formed by C–C–N–C and C–N–C–N coupling as well as dealkylated trimers were found in the solution [14]. N,N-dialkyl anilines are more easily oxidized and the oxidation potentials shift to more positive values in the order of aniline < N,N-di-n-butyl < N,N-diethyl < N-methyl-N-ethyl < N, N-dimethylaniline. Two one-electron oxidation peaks were observed for these dialkyl anilines

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1125

Aromatic Nitrogen–Containing Compounds

in acidic solution of pH = 1 and both products, radical cations and dications, were consumed in follow-up reactions making the observation of cathodic peaks on the reverse scan very difficult [17]. However, only one peak for N,N-dimethylaniline (XX) was observed without any cathodic response in alkaline media of pH = 13, where the radical cation is less stable [18]. The benzidine derivative was confirmed as the main product of the oxidation of XX. Side reactions of the primary radical cation were investigated using the electrochemical thermospray MS technique [19]. Reactions of the neutral radical formed by the deprotonation at the α-carbon atom of the alkyl group were considered in order to explain the formation of dimers with different structures. On the other hand, in acetonitrile solutions the radical cation formed in the one-electron oxidation of XX is so stable that its reduction peak was observed in CV curves recorded at scan rates above 500 V s−1 [20]. The fast radical cation–radical cation dimerization and deprotonation result in the formation of N,N,N′,N′-tetramethylbenzidine as the final dimer, but no evidence of further polymerization was observed. The presence of ring substituents in the para-position in N,N-dimethylanilines can affect radical cation coupling reactions. Radical cations of 4-bromo-N,N-dimethylaniline generated by anodic oxidation in MeCN in the absence of parent molecules form tetramethylbenzidine dication in radical cation–radical cation dimerization, but if parent molecules are present in the solution they dimerize with the primary radical cation yielding tetramethylbenzidine radical cation [21]. In both of the cases, C–C coupling occurs with C–Br bond cleavage. However, the oxidation of N,N-dimethylp-toluidine (XXI) in similar conditions gives N,N,N′,N′-tetramethyl-α,α′-bi-p-toluidine (XXIII) (Scheme 29.4) and rate constants for the formation of the neutral radical XXII· and the final product XXIII were obtained by the simulation of CV curves [22]. The oxidation of diphenylamine (XXIV) in MeCN solution at platinum ultramicro electrodes gives the radical cation XXIV+ • (Scheme 29.5), which can be detected in CV curves recorded at a scan rate above 100 V s−1 [23]. Finally, N,N′-diphenylbenzidine (XXV) formed after the coupling of two radical cations XXIV+ • is further oxidized to a dication. Even though different CV curves were observed after the addition of 2,6-lutidine acting as a weak base, yet the same product XXV was formed [24]. Stronger bases change the reaction pathway probably due to the deprotonation at the nitrogen atom. A similar coupling is impossible in the presence of para-methoxy groups in dianisylamine and the oxidation pathway is more complicated, but in this case not all oxidation products could be identified [24]. Nevertheless, the oxidation at lower scan rates in the presence of lutidine probably Me

+

Me

Me

Me

N

Me N–

(XXI)

Me

Me N

+

CH2

Me (XXI+ )

H–+

N + XXI

–e

Me

Me

Me

(XXII ) + XXII

Me

H

H

Me

C H

C H

N

Me N Me

(XXIII)

SCHEME 29.4

Pathways for the anodic oxidation of N,N-dimethyl-p-toluidine XXI in MeCN.

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1126

Organic Electrochemistry +

H N

H N

–e

(XXIV+ )

(XXIV)

–2H+ +XXIV+ N H

N H (XXV) –e

+

N H

N H –e

SCHEME 29.5

+

+

N H

N H

Pathways for the electrochemical oxidation of diphenylamine XXIV in MeCN.

yields the cyclic product 2,6-dimethoxy-9,10-dianisyl-9,10-dihydrophenazine, also suggested in earlier works. A similar cyclic product 2,6-dimethyl-9,10-diphenyl-9,10-dihydrophenazine is characteristic of the oxidation of 3-methyl-diphenylamine and other 3-substituted and 3,3′-disubstituted diphenylamines [25]. The aromatic substitution of diphenylamines with the formation of cyanodiphenylamines was elaborated by the electrolysis in methanol solutions containing cyanide ions [26]. Two-electron oxidation of p-aminodiphenylamine (VII) observed in MeCN was explained by a mechanism involving the proton transfer between the parent molecule and its radical cation (the rate-determining step [rds]) and the immediate oxidation at the same potential of the neutral radical to the final cation as is shown in Scheme 29.6 [27]. The N–N coupling of the aminyl radical of p-monosubstituted diphenylamines in MeCN under basic conditions results in the formation of corresponding N,N,N′-triaryl p-phenylenediamines and tetraarylhydrazines (p–XC6H4)2N–N(p–XC6H4)2 [28]. Among triarylamines the electrode oxidation in MeCN of triphenylamine (XXVI) to tetraphenylbenzidine (XXVII) was most intensively investigated (Scheme 29.7). Radical cations XXVI+·

–e VII

+VII

VII+

+

N H

NH3

+ +

–e N H

SCHEME 29.6

NH

N H

NH

Mechanism for the anodic oxidation of p-aminodiphenylamine VII in MeCN.

© 2016 by Taylor & Francis Group, LLC

1127

Aromatic Nitrogen–Containing Compounds +

XXVI+

–e N

N

(XXVI)

(XXVI+

N

N

–2 H+ (XXVII) –e

)

XXVII+ –e

+N

N+

(XXVII2+)

SCHEME 29.7

Mechanism for the electrode oxidation of phenylamine XXVI.

formed in the first step are relatively stable (their reversible reduction is observed at 20 V s−1) and the stability increases for derivatives with para substituents when the deprotonation necessary for the dimerization process is difficult [29]. The radical cation dimerization according to the DIM1 mechanism is the rds, and the dimeric product XXVII is consecutively oxidizable to the radical cation and the quinoidal dication at less positive potentials than it was produced [30]. It can be added that a number of tri-p-substituted triarylamines representing a wide range of their oxidation potentials and giving stable radical cations were successfully used as mediators in indirect electrochemical oxidation of different compounds [31] (see also Chapter 15).

C. RING-SUBSTITUTED ANILINES In aqueous acidic media, para-substituted anilines (XXVIII) are oxidized to radical cations that undergo rapid head-to-tail coupling with the formation of protonated 4′-substituted 4-aminodiphenylamine (XXIX). The overall process given in Equation 29.1 corresponds to a oneelectron transfer for one aniline molecule if the leaving para-substituent is in the anionic form (e.g., halide or methoxide ion), but it corresponds to a two-electron process if a neutral group is left (e.g., CO2 for p-aminobenzoic acid) and the final dimer occurs in the reduced form [32]. 2 p-X − C6H 4 − NH 2 − 2e → p-X − C6H 4 − N=C6H 4 =N + H 2 + X – + 2H + ( XXVIII )

(29.1)

( XXIX )

However, for the oxidation of p-toluidine in 0.1 M H2SO4 (and in MeCN), CV investigations showed the head-to-head coupling of radical cations with the formation of 4,4′-dimethylhydrazobenzene, which is irreversibly oxidized at the potential of formation to the main product 4,4′-dimethylazobenzene [8]. In solutions with pH < 8, another product 2-amino-4′,5-dimethyldiphenylamine formed in the ortho-coupling was proposed earlier [33]. A similar overall reaction as given by Equation 29.1 was found in aprotic media (DMF and MeCN); however, the product may or may not be protonated depending on the medium basicity [34–37]. Moreover, during the oxidation of p-haloanilines in aprotic media, the leaving anions Cl− or Br− are oxidized to dihalogens (Equation 29.2), which form with parent anilines 2,4-dihaloanilines as minor

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1128

Organic Electrochemistry +

NH2 + XXVIIIa MeO

–e

OMe

H + NH

–

NH2

OMe

OMe NH2

–MeOH

(XXVIIIa)

+

H N

MeO

NH2 –e

MeO

SCHEME 29.8

+

+

N H

NH2

Detailed mechanism for the electrooxidation of p-anisidine XXVIIIa in DMF.

products, as will be discussed further on [35–37]. Thus, depending on the fraction of the oxidized halide ions the apparent number of electrons per aniline molecule changes from 1 to 1.5. 2 p-X − C6H 4 − NH 2 − 3e → p-X − C6H 4 − N =C6H 4 = NH + ½ X 2 + 3H + ( XXVIIIb ) X = Cl ( XXVIIIc ) X = Br

(29.2)

( XXX )

At high scan rates using gold ultramicroelectrodes, reversible one-electron CV peaks were found in DMF for the oxidation of p-anisidine (XXVIIIa), p-chloroaniline (XXVIIIb), and p-bromoaniline (XXVIIIc) supporting the formation of corresponding radical cations in the first step [34,35]. However, the mechanism of their decay is not the same for all aforementioned reactants, as indicated by different slopes of the change of anodic peak potentials with the scan rate ∂E pa /∂logν and with the reactant concentration ∂Epa /∂logc. For XXVIIIa (X = OMe), both the aforementioned slopes are equal to 30 mV per decade, indicating a dimerization according to the DIM2 mechanism as the second-order rds. The detailed mechanism shown in Scheme 29.8 involves the nucleophilic attack of parent anisidine XXVIIIa onto its radical cation (the rate constant of this rds is 7.7 M−1 s−1), further rearrangement of a radical, the elimination of methanol, and the oxidation to the final dication [34]. It should be added that a more complicated mechanism of the decay of radical cations of XXVIIIa was found in MeCN using the electron-transfer stoppedflow method with spectral detection of intermediates that however are not formed in the vicinity of the electrode [38]. On the other hand, for the oxidation of p-haloanilines, XXVIIIb and XXVIIIc in DMF slopes of ∂E pa /∂logν = 20 mV per decade were found [35], but peak potentials were not dependent on the reactant concentration. Moreover, the addition of 2,6-lutidine at a higher concentration changes the slope of ∂Epa /∂logν to 30 mV per decade, indicating the irreversible reaction with the base. Thus, the mechanism proposed (Scheme 29.9) involves a fast reversible deprotonation of the primary radical cation by any base present in the solution with the formation of the neutral radical XXXI•, followed by its coupling with the primary radical cation (the rds) and spontaneous elimination of HX giving the final 4-imino-4′-halodiphenylamine [35]. Note that no similar deprotonation occurs for p-anisidine radical cation XXVIIIa+ •, a much weaker acid, which explains the different DIM2 mechanism observed. The controlled potential electrolysis of p-chloroaniline (XXVIIIb) and p-bromoaniline (XXVIIIc) in MeCN affords a number of products besides halodiphenylamines XXX shown in Equation 29.2. The product distribution was reported as the relative intensity of peaks in gas

© 2016 by Taylor & Francis Group, LLC

1129

Aromatic Nitrogen–Containing Compounds +

NH2

NH2

Base

BaseH+

NH

NH

X

X

–e

X

X

(XXXI ) (XXVIIIb) X = Cl (XXVIIIc) X = Br NH

–

NH

+

N

–

N H

H

–HX

+

X

X

H X

SCHEME 29.9 Proposed mechanism for the anodic oxidation of p-haloanilines XXVIIIb and XXVIIIc in DMF.

chromatography-electrospray MS measurements [36,37]. For XXVIIIb, 4-amino-4′-chlorodiphenylamine (78.4%) was the main product formed in the head-to-tail coupling, but the coupling in the ortho-position yielded 2-amino-4,5′-dichlorodiphenylamine (30.4%) and the reaction with elemental chlorine mentioned earlier gave 2,4-dichloroaniline (23%). On the other hand, for XXVIIIc the main products were 2,4-dibromoaniline (58%) and 4-amino-4′-bromo-diphenylamine (38.8%) but ortho-coupling was very small as expected for the larger bromine atom. However, recent voltammetric and in situ ESR study of the oxidation of p-chloro, p-bromo, and p-iodoaniline in MeCN at gold electrodes indicated that no single reaction mechanism alone can explain all results found [39]. Apart from the fast route producing the nonparamagnetic dimer at a higher oxidation state as shown in Scheme 29.9, there is the competitive slow reaction of deprotonated neutral radical XXXI• with the parent aniline resulting in the formation of the unstable dimeric radical cation that can be detected by ESR using a tubular flow cell [39]. The electrochemical behavior of p-haloanilines is quite different in more basic media when head-to-head coupling giving rise to azobenzene derivatives is the main oxidation route, similarly as for unsubstituted aniline [10]. Such behavior was reported for p-chloroaniline (XXVIIIb) in MeCN solutions containing 0.1 M pyridine [40]. In the last case, the main product of the large-scale electrolysis of XXVIIIb was 4,4′-dichloroazobenzene (24.3%) formed by the radical–radical coupling of neutral radical p–Cl–C6H4 –NH• produced in the deprotonation of primary radical cation by pyridine. The two-electron oxidation of p-aminophenol and its derivatives in acidic solutions gives quinoneimine that hydrolyzes to p-quinone [41]. A similar two-electron oxidation to quinonediimine was found for p-phenylenediamine H2N–C6H4 –NH2 (XXXII), but in solutions of pH from 2 to 6, the radical cations formed after the one-electron step are stable and they were detected by ESR measurements [42]. The reversible one-electron oxidation of XXXII yielding stable radical cations was also observed in many aprotic solvents, and this redox couple was used as a model reactant for the elucidation of solvent effects on heterogeneous rate constants [43,44]. Such investigations were then extended to N,N,N′,N′-tetramethyl-p-phenylenediamine [45] and to binary solvent mixtures [46]. On the other hand, head-to-tail and tail-to-tail coupling was observed for the anodic oxidation of ortho- and meta-substituted anilines (XXXIII) in aqueous acidic media resembling the behavior of aniline itself [32]. A detailed mechanism of the oxidation in 3 M H 2SO 4 solutions proposed by Hand and Nelson [47] is shown in Figure 29.1. The primary radical cations formed in the first electron transfer undergo tail-to-tail coupling to benzidines XXXIV or produce diphenylamines XXXV (in particular for molecules with electron-withdrawing ortho-substituents).

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1130

Organic Electrochemistry +

++

NH2

NH2

NH2 + XXXIII+

–e X

H2N X

X

X

(XXXIII)

(XXXIII+ )

H+

NH2 X

(XXXIV)

2e

X NH2

+

+

NH2

H2N

NH

X

X + XXXIII X

H N

NH2 X

X

meta

ortho

(XXXV) 2e + 2H+

NH

NH2

H3O+

+ X

X

Fast

N

NH

X

X

H3O+

N

Fast X

O

O

X

H3O+

Slow

H3O+

Slow

O OH

H2O

X

O

NH2 X

H2O

+

X

X

O

OH

O

OH 2e + 2H+ O X OH (XXXVI)

O

FIgURE 29.1 Reaction pathways proposed for the anodic oxidation of ortho- and meta-substituted anilines in 3 M H2SO4. (Adapted from Hand, R.L. and Nelson, R.F., J. Electrochem. Soc., 125, 1059, 1978. Reproduced by permission of ECS—The Electrochemical Society.)

Further oxidation and hydrolysis occur by different routes for ortho- and meta-substituted reactants, but the same final product, 2-substituted para-benzoquinone XXXVI, is formed. A similar mechanism was also observed in solutions of 2 M H2SO 4 [9], but not in 0.1 M H2SO 4 (and MeCN) where the formation of hydrazobenzene, which is next oxidized to azobenzene, was found for all three isomers of toluidine [9]. On the other hand, as a result of the anodic oxidation of o-aminophenol (XXXVII) in aqueous neutral and alkaline solutions on a silver electrode, the linear dimer 2,2′-dihydroxyazobenzene (XXXVIII) produced by the N–N coupling of two neutral radicals XXXIX was obtained as the main product. However, the cyclic dimer 3-aminophenoxazone (XL) formed by C–N coupling was the dominant product in acidic solutions [48] as is shown in Scheme 29.10. XL can be electrochemically reduced. Cyclic dimers produced by the C–N coupling of radical cations are also proposed [49] for the oxidation of o-anisidine in

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1131

Aromatic Nitrogen–Containing Compounds

+

NH2

NH2 OH

OH

–e

NH

+ OH– –H2O

+XXXIX N N

–2e, –2H+ OH

OH HO (XXXVIII)

(XXXIX) +XXXVII+

(XXXVII+ )

(XXXVII)

N

NH2

O

O

(XL) +2e +2H+

SCHEME 29.10

H N

NH2

O

OH

Pathways for the anodic oxidation of o-aminophenol XXXVII in aqueous solutions.

acidic solutions on a platinum electrode. The aforementioned conclusions were obtained [48,49] using surface enhanced Raman scattering spectra. More recently, absorption spectra of the oxidation products of ortho- and meta- isomers of toluidine and anisidine in MeCN were investigated using the electron-transfer stopped-flow technique [50]. The formation of 3,3′-dimethylbenzidine (in the form of radical cation and dication) by the C–C coupling in para-position was proven for the oxidation of o–toluidine. Consecutive reactions of that dimer with monomer radical cations were reported. However, a similar Cpara –Cpara coupling of o–anisidine radical cations was excluded, and moreover, for more reactive radical cations of m-methyl and m-methoxy substituted anilines, the spectra could not be observed at all [50]. Among polysubstituted anilines, the anodic behavior of halogen derivatives in MeCN and in aqueous solutions was investigated most intensively. The oxidation of 2,4-dihaloanilines (XLI) in MeCN showed a few different products and the overall reactions for their formation are given in Scheme 29.11 [36,37,51]. The main products are N-(2,4-X2–phenyl)-3′-X-quinone diimine (XLII) X +XLI –4e, –4H+

X N

X

NH

X (XLIV) X = Cl, Br NH2

X X

+2 XLI –4e, –4H+

X N

X

X (XLIa) X = Cl (XLIb) X = Br

Cl Cl

X

NH + X (XLIIIa) X = Cl (XLIIIb) X = Br

(XLIIa) X = Cl (XLIIb) X = Br

+3 XLIa –8e, –8H+

– NH2

X

Cl NH + XLIIa

N Cl

Cl (XLVa)

SCHEME 29.11

Formation of different products during the oxidation of 2,4-dihaloanilines XLI in MeCN.

© 2016 by Taylor & Francis Group, LLC

1132

Organic Electrochemistry Br 2 XLIIIb

–3e, –2 H+ –1/2 Br2

Br

Br N

NH

Br

Br (XLVb)

Cl

Cl –e XLIIIa

XLIIIa+

+XLIIIa+ –2H+

H H N N

Cl Cl

Cl Cl

(XLVI)

SCHEME 29.12

Anodic oxidation of 2,4,6-trihaloanilines XLIII in MeCN.

and trihalogenated aniline (XLIII). The formation of XLII in the head-to-tail coupling of radical cations (and the elimination of halogen) is similar to the mechanism found by Bacon and Adams [32] for para-substituted anilines in acidic media. Indeed in fast measurements using a rotating disk electrode, the one-electron transfer corresponding to the oxidation of one XLI molecule was found [51]. However, under the controlled potential electrolysis two-electron processes (for one reactant molecule) give halogenated by-products XLIV and XLVa [36,37]. On the other hand, for the oxidation of 2,4,6-trihaloanilines (XLIII) in MeCN the apparent number of electrons equal to 1.5 and additional CV peaks indicated the oxidation of halide ions expelled from the para-position [36,37] similarly as it was observed for p-haloanilines. Thus, the main product for the oxidation of XLIIIb is N-(2,4,6-Br3 –phenyl)-3′,5′-Br2-quinone diimine (XLVb) as is shown in Scheme 29.12 [37]. However, XLVa is a minor product for the oxidation of chloro-derivative XLIIIa, but the main product 2,2′,4,4′,6,6′-hexachlorohydrazobenzene (XLVI) is formed by the head-to-head coupling of radical cations in the reaction shown also in Scheme 29.12 [36]. In the presence of one or two methyl groups (which are electron-releasing) in ortho-position, a stabilization of radical cations is observed. Therefore, two one-electron CV anodic peaks corresponding to the consecutive formation of radical cations and dications were observed for the oxidation of 2,4and 2,6-dimethylanilines in 0.1 M H2SO4 [9]. In the case of 2,4-derivative, the ensuing dimerization produces substituted hydrazobenzene and, after its instantaneous oxidation, azobenzene. However, for the 2,6-derivative the para-position is vacant and all three types of coupling occur similarly as it was reported for the aniline itself (cf. Scheme 29.2) [9]. On the other hand, in ACN the anodic behavior of 2,4- and 2,6-dimethylanilines resembles that of p-toluidine and o-toluidine, respectively [9]. Anodic behavior of polyhaloanilines at platinum electrodes in aqueous media with higher acidity (at least 1.0 M H 2SO 4) was quite different because the reactants are completely protonated and no dimerization is observed but a hydrolysis of primary intermediates [52–55]. The two-electron oxidation of protonated 2,4-dihaloanilines (XLVII), which is the rds, gives the intermediate tripositive cation XLVIII that loses a proton (Scheme 29.13), and next by hydrolysis, it gives two products: 2-halo-p-benzoquinoneimine (XLIX) and 2-halo-p-benzoquinone (L). Both these products can be reduced giving rise to characteristic CV peaks. The intermediate cation XLVIII can also lose two protons from the benzene ring producing the electroactive species responsible for the different oxidation process observed at higher acidity but giving the same final products XLIX and L. Both these processes of the oxidation of XLVII are controlled either by diffusion or by adsorption depending on the medium acidity [53,54]. Moreover, after prolonged electrolysis of the bromo-derivative XLVIIb, a slow hydration of bromobenzoquinone Lb produces 3-bromo-1,2,4-benzenetriol (LI), which is reversibly oxidized as shown in Scheme 29.13 [53].

© 2016 by Taylor & Francis Group, LLC

1133

Aromatic Nitrogen–Containing Compounds +

+

++

NH3

NH3

+

NH2

NH2

X

X

–2e

+ +

X

–H+

(XLVIIa) X = Cl (XLVIIb) X = Br

+H2O

X

+H2O –NH4+

–2H+, –X– X

X

O X

O

O

X

(XLVIII)

(L)

(XLIX) +2e +2H+

O

+

O

OH Br

Br

+H2O

+2e +2H+ OH

NH3 Br

X

–2e

X

+

OH O

OH

(Lb)

(LI)

–2H

OH O

OH

OH

SCHEME 29.13 Mechanism for the anodic oxidation of protonated 2,4-dihaloanilines XLVII in acidic media.

The anodic oxidation in acidic media of 2,4,6-trisubstituted anilines with the Br atom in paraposition depends on the nature of the ortho-substituents. For weak electron-withdrawing Br atoms, a two-electron irreversible process is followed by deprotonation and further hydrolysis. Thus, 2,4,6-tribromoaniline in its protonated form LII, which predominates in strong acidic media, is oxidized at platinum electrodes [52] with the formation of tripositive cation LIII (Scheme 29.14) similarly as it was found for the oxidation of 2,4-dibromoaniline XLVIIb. The next loss of a proton can occur at different positions giving two possible cations LIV and LV. The former is hydrolyzed first to 2,6-dibromobenzoquinone (LVI), which is not electroactive in CV experiments. However, its further hydrolysis yields 3,5-dibromo-1,2,4-trihydroxybenzene (LVII), which can be reversibly oxidized (Scheme 29.14) [52]. The latter cation LV is oxidized in a two-electron process, but its dimerization with the parent unprotonated aniline XLIIIb was also suggested [52]. On the other hand, in the presence of a stronger electron-withdrawing nitro substituent in the ortho-position, the radical cation formed after the first electron transfer is more stable. Therefore, 2,4-dibromo-6-nitroaniline in solutions of sulfuric acid (where the amino group is protonated) is +

+

+

NH3

NH3

NH2

Br

Br

–2e

Br

Br + +

–H+

O Br

Br

+2 H2O +

–HBr, –H

+

Br

Br

, –NH4+

Br

Br

Br

O

(LII)

(LIII)

(LIV)

(LVI)

–H+

+H2O

+

NH3

O Br

Br

Br

–2e

+

OH Br (LV)

SCHEME 29.14

OH Br

O

Br

Br

–2H+

OH OH (LVII)

Pathways for the anodic oxidation of 2,4,6-tribromoaniline LII in strong acidic media.

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1134

Organic Electrochemistry +

+

NH3

NH2 X

Me

–e

LVIII2+

–H+

O X

Me

Br

Me

+ LVIIIa Br

Br

(LVIIIa) X = Br (LVIIIb) X = Me

(LIX + )

Br O (LXII) +H2O, Br2 –Br–, –H+

+LIX+ –Br–, –H+ Me

Me +

NH+2

N H

Br

Me

X

+H2O –NH4+

SCHEME 29.15

O

X

X

X

(LXIa) X = Br (LXIb) X = Me +2e, +2H+

+2e, +2H+

Br

N H

Br

(LXa) X = Br (LXb) X = Me

Me H2 N + X

Me +

Me NH3+ X

Br

Me H2 N + X

Me OH X

Mechanism for the anodic oxidation of protonated o-substituted p-bromoanilines LVIII.

oxidized according to the eCe mechanism where two reversible one-electron steps are divided by the irreversible chemical reaction of the loss of one proton (the rds) [55]. Further hydrolysis of the resulting dipositive cation (similar to LIV but with the nitro group instead of one ortho-Br atom) with the loss of Br− and NH4+ ions gives rise to 2-bromo-6-nitro-p-benzoquinone that can be reversibly reduced. Thus, the aforementioned oxidation process is similar to that proposed by Hand and Nelson for ortho- and meta-monosubstituted anilines [47]. Finally, the presence of one or two electron-releasing o-methyl groups in p-bromoaniline stabilizes radical cations due to the inductive effect of o-substituents changing the oxidation in acidic media to the eC mechanism. It includes (Scheme 29.15) the reversible electron transfer from the protonated reactant LVIII and the deprotonation of the formed dication radical LVIII 2+ • to radical cation LIX+ • (the rds). The head-to-tail coupling of two LIX+ • gives finally diimine LX, which is hydrolyzed to benzoquinoneimine LXI, and both these products can be further reduced in CV experiments as is shown in Scheme 29.15 [56]. Therefore, the final product of the controlled potential electrolysis of LVIIIb (X = Me) is LXIb. However, for only one ortho-methyl group in LVIIIa (X = Br) a reactant has less electron-releasing character and further hydrolysis and deprotonation of LXIa occurs followed by bromination. This gives 2-methyl-5,6-dibromop-benzoquinone (LXII) as the final product that can be further reduced to the corresponding hydroquinone [56]. Extremely stable intermediates—nitrenium ions, nitryl radicals, and dications—can be formed in the anodic oxidation of sterically hindered anilines, in particular 2,6-di-tert-butyl-4-(4′-phenyl)anilines (or 4-amino-3,5-di-t-butyl-biphenyls, Scheme 29.16) investigated by Speiser and coworkers [57–62]. The formation of radical cations LXIIIa+ •, LXIIIb+ •, and LXIIIc+ • coloring MeCN solutions to deep blue, green, and orange, respectively, during the reversible anodic oxidation of the corresponding anilines LXIII was confirmed by electroanalytical and ESR measurements as well as by UV/Vis spectra obtained using modulated specular reflectance spectroscopy (MSRS) [57]. No indication of follow-up reactions of these radical cations was found in the experimental time scale.

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1135

Aromatic Nitrogen–Containing Compounds

NH2

R-C6H5NH3+

–e

(LXV+)

–H+ –e

R-C6H5NH22+

R-C6H5NH2

R-C6H5NH2+

(LXIII)

(LXIII+ )

(LXIII2+ )

+H+ –H+

+H+ –H+

+H+ –H+

R-C6H5NH– (LXIV–)

SCHEME 29.16 in MeCN.

+H+ –e

(LXIIIa) R = Ph (LXIIIb) R = p-C 6H4OMe (LXIIIc) R = p-C 6H4NMe2 (LXIIId) R = p-C 6H4NO2 (LXIIIe) R = t-Bu

+

(LXV++ )

+H+ –H+ R

R-C6H5NH3+

–e

–e R-C6H5N H (LXIV )

R-C6H5NH+ (LXIV+)

Detailed pathways for the anodic oxidation of a series of sterically hindered anilines LXIII

Their stability was attributed to the electron-donating effect of 4-phenyl substitution, which is also responsible for the shift of standard potentials to low values and the low wavelength bands to lower energy. The observed decay of the radical cation LXIIIe+ • (with a 4-t-butyl group instead of phenyl) confirmed the aforementioned explanation [57]. For a series of 18 anilines LXIII with different substituents, the formal potentials of the reversible one-electron oxidation were correlated with Hammett σ constants and Taft’s dual substituent parameters [59]. The results showed that the Gibbs free energy change depends mainly on the electronic structure of radical cations LXIII+ • in which two phenyl rings have a coplanar conformation allowing strong π-electron interactions and the stabilization of the positive charge. On the other hand, in neutral anilines LXIII a steric twisting between two phenyl rings prevents the conjugation. For electron-withdrawing substituents (e.g., –NO2), an additional cathodic peak is observed probably corresponding to the reduction of the neutral radical LXIV •. Moreover, it was found that the oxidation mechanism of LXIIIc with two reaction sites seems to be different; that is why the oxidation of the –NMe2 group instead of the –NH2 one was considered. The rate constants for the protonation of LXIII to anilinium cations LXV+ in MeCN containing perchloric acid and the back deprotonation determined from voltammograms by the multiparameter estimation [60] were considerably lower than expected for diffusion-controlled reactions supporting the effects of the steric hindrance in ortho positions. Consequently, anilines LXIII are less basic than the unsubstituted aniline: for the anilinium cation XIII pKa = 10.56 in MeCN whereas pKa values of cations LXV+ are much smaller, ranging from 3.70 for LXVc+ to 4.61 for LXVa+ and LXVb+ [60]. Further oxidation of radical cations leads to dications (Scheme 29.16), but only LXIIIc2+ is stable and a simulation of CV curves supported the ee mechanism for the oxidation of LXIIIc in two oneelectron steps. Dication LXIIIb2+ decays moderately fast, but the oxidation mechanism of LXIIIb is complicated. It includes the deprotonation of LXIIIb2+ to nitrenium ion LXIVb+ (which can be reduced to nitryl radical LXIVc• giving an additional CV peak) and the protonation of parent aniline yielding cation LXVb+, which is oxidized at potentials close to the oxidation of LXIIIb+•. UV/Vis spectra of dications LXIIIc2+ and LXIIIb2+ were obtained using MSRS technique. On the other hand, dication LXIIIa2+ decays extremely fast and could not be detected [58]. Moreover, the anodic oxidation of LXIIIc in MeCN with an excess of 2,6-lutidine resulted in the formation of persistent nitrenium ion LXIVc+ giving a deep purple color of the solution. The last ion was identified on the basis of CV experiments as well as UV and 1H NMR spectra. NMR signals confirmed the delocalization of the positive charge over both rings of the biphenyl moiety [61]. Reactions of electrogenerated nitrenium ions LXIVa+, LXIVb+, and LXIVc+ with a number of nucleophiles were also reported [62].

© 2016 by Taylor & Francis Group, LLC

1136

II.

Organic Electrochemistry

OTHER AROMATIC AMINES

In general, the anodic behavior of polynuclear aromatic amines, for example, the derivatives of anthracene and fluorene as well as the more recently studied 1-naphthylamine LXVI [63,64], 1-pyrenamine [65] and derivatives of dibenzothiophene [66], is similar to that observed for anilines. In particular, the oxidation of LXVI in MeCN corresponds to the eCe mechanism with the reversible formation of radical cations LXVI+ • in the first step and their very fast follow-up reactions forming the dimer, which is easier to oxidize in a two-electron process than the parent amine. Among three possible paths of dimerization (similarly as shown in Scheme 29.2 for aniline), the main product 4-amino-1,1′-dinaphthylamine is formed by the head-to-tail (i.e., C4 –N) coupling. Moreover, the effects of the addition of an acid and a base suggested that the deprotonation of radical cations LXVI+ • to neutral radicals precedes their dimerization. However, naphthidine, the product of the tail-to-tail (i.e., C4 –C4′) coupling was also identified in CV curves [63]. On the other hand, in the more basic solvent DMSO the main oxidation product 1,1′-hydrazonaphthalene is produced in the head-to-head (i.e., N–N) coupling, but naphthidine was not found among the products [64]. A  similar eCe mechanism was proposed for the oxidation of 1-pyrenamine in MeCN [65]. The results of the double potential step chronoamperometry indicated the coupling of two radical cations (but not the radical cation with the parent amine) and made it possible to determine the corresponding rate constant [65]. The reversible one-electron oxidation of p-phenylenediamine XXXII in aprotic solvents was already mentioned in Section I.C, and many earlier investigations of their derivatives were reviewed [3]. The slow two-electron oxidation of 3,6-bis(dimethylamino)durene and 9,10-bis(dimethylamino) anthracene in MeCN [67] was ascribed to large inner reorganization energy due to substantial structural changes. On the other hand, for the anodic oxidation of amino derivatives of dibenzothiophenes in MeCN containing LiClO4, the reversible formation of stable radical cations was shown, and the most stable of them obtained from 3,7-diamino and 3,7-bis(dimethylamino)dibenzothiophene were isolated in the form of crystalline perchlorate salts and were characterized by ESR spectra [66]. In recent years, the anodic behavior of polyaminobenzenes with an increasing number of –NH2 groups was also intensively investigated [68–73]. High-spin organic cations, inspiring for the preparation of magnetic materials, were obtained [68] in the anodic oxidation of N,N,N′,N′,N″,N″hexaanisyl-1,3,5-triaminobenzene in butyronitrile solution. Two reversible one-electron steps at room temperature and one more step reversible at −78°C were observed in CV curves at 0.2 V s−1. These results indicated the consecutive formation of radical cation, dication and trication. ESR spectra of those ions obtained by chemical oxidation showed the doublet, triplet, and quartet spin state, respectively [68]. First oxidation potentials of similar 1,3,5-tris(diarylamino)benzenes with other substituents were also reported [69]. An early study of the newly synthesized 1,2,4,5-tetrakis(dimethylamino)benzene LXVIIa (Scheme 29.17) showed [70] that the first oxidation step in MeCN is a two-electron process producing the structurally distorted dication with no through-conjugation in the molecule. That was contrary to the successive one-electron oxidations of N,N,N′,N′-tetramethyl-p-phenylenediamine in Me2N R Me2N

NMe2 R

LXVIIa

–e

LXVIIa+

–e

LXVIIa2+

in CH2Cl2 + Bu4N[B{C6H3(CF3)2}4]

NMe2

(LXVIIa) R = H (LXVIIb) R = NMe2

SCHEME 29.17

–2e LXVIIb LXVIIb2+ in MeCN/CH2Cl2 + Bu4NPF6

–e

LXVIIb3+

–e

LXVIIb4+

–2e

Anodic oxidation of tetrakis and hexakis(dimethylamino)benzenes LXVII.

© 2016 by Taylor & Francis Group, LLC

LXVIIb6+

Aromatic Nitrogen–Containing Compounds

1137

MeCN [71] and hexaaminobenzene in nitromethane [72] producing radical cations in the first step. However, in a recent careful analysis of the anodic properties of LXVIIa and similar compounds in different media, Adams et al. [73] showed that the difference between the first and the second oxidation potentials depends not only on structural changes during these processes but also on the solvent and the electrolyte used. In particular, the stability of dications is enhanced by their strong solvation and ion pairing with electrolyte anions. Both these phenomena favor the disproportionation of radical cations and thus the two-electron step. Consequently, it was shown in an elegant way that two one-electron CV steps could be obtained for the oxidation of LXVIIa in the nonpolar solvent CH2Cl2 containing as a supporting electrolyte [Bu4N][B{C6H3(CF3)2}4] with large and noncoordinating anions (Scheme 29.17). Moreover, crystal salts of radical cations LXVIIa+ • and dications LXVIIa2+ (and a few similar compounds) were isolated and their structures determined by x-ray diffraction [73]. The electrochemical oxidation of hexakis(dimethylamino)benzene LXVIIb in an MeCN/ CH2Cl2 mixture (1:1, v/v) containing NBu4PF6 showed in the first step a two-electron electrochemically quasi-reversible process [74,75], whereas the one-electron reversible process was reported earlier in CH2Cl2 at −50°C [76]. Both results look consistent with the aforementioned findings of Adams et al. [73]. The slow two-electron kinetics [74] was related to the significant distortion of the planar aromatic ring during the oxidation. Namely, the stable dication formed has two noncoplanar polymethine units carrying one positive charge in each and connected by two single C–C bonds. Such a structure was confirmed by X-ray examination of the salt LXVIIb2+ (PF6 −)2. Further oxidation produced tri-, tetra-, and hexacations (Scheme 29.17), which undergo slow follow-up reactions [74,75]. Computer simulation of CV curves at different scan rates made it possible to estimate standard potentials and rate constants for those electrode and chemical reactions [74,75]. The increasingly greater sluggishness of electron transfers related to greater structural changes was observed for the two-electron oxidation of monomeric peralkylated hexaamino(1,3)metacyclophane in the same medium showing an unusually big separation of anodic and cathodic peaks [77]. On the other hand, sterically rigid dodecaamino(1,3,5)-cyclophane with almost planar and parallel two hexaaminobenzene units (for which the twisted structure in a higher oxidation state is not possible) showed two reversible one-electron oxidation steps at close potentials [77]. CV investigations of a number of meta-connected oligoarylamines showing a number of consecutive oxidation products that are stable in aprotic solvents were recently overviewed [1]. Among synthetic macromolecules, the redox-active dendrimers gained increased attention over the last two decades. Highly symmetrical hexaarylbenzene derivatives with six triarylamine redox moieties gave two reversible pairs of CV peaks during the overall six electrons oxidation in CH2Cl2 [78]. Moreover, the selectively generated radical trication showed a strong intervalence charge-transfer band in the near infrared (NIR). Potentials of the one-electron oxidation of polyamidoamine-type dendrimers with p-phenylenediamine redox core (reversible in methanol but slow in some aprotic solvents) obtained from the differential pulse voltammetry decrease with increasing size (and the related electronic nature) of low-generation dendrimer chains [79]. The oxidation process strongly affects the preferred geometry of chains, but not that of the dendrimer core [79]. Electrochemically generated radical cations were characterized by ESR and UV/Vis NIR spectra [80]. References to earlier works can be found in Reference 79.

III.

AZO AND AZOXy COMPOUNDS

Azo compounds can be electrochemically reduced in aprotic as well as protic media and some substituted derivatives can also be oxidized. The cathodic reduction of azobenzene (XVI), the most intensively investigated representative, in MeCN and DMF containing tetraalkylammonium salts where ion pairing is absent, occurs at mercury electrodes in two one-electron steps (Scheme 29.18). The first of them is reversible and produces relatively stable radical anions XVI− • detectable by the ESR method [81–83]. The second electron transfer is generally also

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1138

Organic Electrochemistry

Ph N

N Ph (XVI)

+e

[ Ph N

N Ph ] –

+e

[ Ph N

N

Ph ]2 –

fast +SH, –S–

(XVI– )

[ Ph N

NH

Ph ] –

(LXVIII)

SCHEME 29.18

Electroreduction of azobenzene XVI in aprotic solvents.

reversible, but it yields more basic dianions that are protonated in a fast chemical reaction with any proton donors, SH, available in solutions [82–84]. For example, the protonation of XVI2− ion in DMF by residual water is reversible and the pKa of the formed PhNHN−Ph is 38.1 [85]. The chemically reversible second electron transfer to XVI in DMF containing neutral activated alumina was observed after cooling [86]. However, the corresponding anodic peak disappears after the addition of MeCN due to its deprotonation by azobenzene dianions [86]. Thus, the observed irreversibility of the second CV peak depends on the solvent and the sweep rate used. The presence of substituents decreasing the basicity of the dianion (as e.g., the para-nitro group) also favors the reversibility of the second step [87]. Products of both reduction steps are oxidized in air to the parent azo compound. Slow decomposition of the dianion or the protonated dianion to arylhydrazine in DMF was reported [82] but in the presence of an excess of proton donors hydrazobenzene (VIII) was the final product. Note that the electrolytic generation of dianions of substituted azobenzene and other aromatic azo compounds for the preparative use as strong bases is reviewed in Chapter 43. 2e LXVIII + OH − −  → XVI + H 2O

(29.3)

XVI + LXVIII + OH − ⇌ 2XVI −· + H 2O

(29.4)

Radical anions XVI− • (and analogous reactants bearing some substituents) formed in the first reduction step in aprotic media can disproportionate to the parent neutral molecules and dianions. In particular, this reaction occurs in the presence of alkali metal cations used in a supporting electrolyte due to ion pairing phenomena [88]. The complicated disproportionation kinetics of ion pairs formed by XVI− • with lithium cations in DMF were studied and the corresponding rate constants were estimated [89]. A rapid disproportionation of protonated radical anions of benzo[c]cinnoline generated by the electroreduction in tetrahydrofuran containing benzoic acid was also investigated [90]. The effects of added proton donors of different acidity on polarographic curves of azobenzenes were reported by Boto and Thomas (cf., review [87]), whereas the effects on CV curves in DMF by Cheng and Hawley [85] and more recently by Astudillo Sánchez and Evans [83] in MeCN. The formation of the monoanion of hydrazobenzene LXVIII shown in Scheme 29.18 was confirmed [83] by the additional anodic peak observed after the addition of water. This peak corresponded to the oxidation of LXVIII in reaction (29.3), whereas it was absent without added water because LXVIII vanishes in reaction (29.4), unless it is eliminated at scan rates higher than 10 V s−1. On the other hand, after the addition of water, hydrogen-bonded complexes XVI− • (H2O) and XVI− • (H2O)2 were formed by radical anions [83]. Their formation constants equal to 3 and 1 M−1, respectively were found from the shift of formal potentials with the water concentration. However, the kinetic analysis also indicated the formation of a very weak XVI− • (H2O)3 complex. The proton transfer from water molecule to radical anion within that complex was proposed in a detailed reaction mechanism [83]. It is followed by the electron transfer from the electrode (or in solution from XVI− •), finally resulting in the formation of hydrated LXVIII (H2O)2.

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1139

Aromatic Nitrogen–Containing Compounds Strongly alkaline media Ph N N Ph

+2e

Ph N

+2H2O

N Ph2

–2 OH–

Ph NH NH Ph (VIII)

(XVI) +H+ Ph N NH Ph

+e

+

SCHEME 29.19

+e

Acidic media Ph N NH Ph

+H+

Ph NH NH Ph +

Electrode reduction and oxidation processes of azobenzene XVI in aqueous media. +H+ VIII

+

Ph NH2 NH Ph

+H+

+

+2e

+

Ph NH2 NH2 Ph

2 Ph NH2 (I)

+

+

H3N-C6H4-C6H4-NH3

SCHEME 29.20

Formation of aniline I by further reactions of hydrazobenzene VIII in acidic solutions.

The reversible reduction–oxidation processes of azobenzene in protic media, in particular the most widely investigated water/alcohol systems, correspond to equilibria of the azo-hydrazo redox couple. They depend strongly on the medium used (first of all on its acidity, but also on the nature of a nonaqueous component) and the concentration of the reactant because of its adsorption. The reaction sequences in acidic and strongly alkaline solutions are shown separately in Scheme 29.19 [2,4]. In both media, linear dependences of reduction potentials on pH were observed; for example, the results obtained in 50% aqueous methanol using a pyrolytic graphite electrode were compared [91] with literature data obtained at mercury electrodes. Hydrazobenzene (VIII) is the same final product of the reduction in protic media of different pH. However, in acidic solutions (pH < 4) instead of a two-electron reversible reduction to VIII an overall four-electron irreversible process occurs including protonation of VIII formed in the first electrochemical step (Scheme 29.20) and its cathodic reduction in the next step with N–N bond cleavage and the formation of aniline (I) [92]. It was shown [92] that unprotonated VIII is not active at electrodes and monoprotonated cation catalyzes the hydrogen evolution. Only diprotonated hydrazobenzene undergoes reduction, but its rearrangement to diprotonated benzidine that adsorbs at electrodes can additionally complicate the process at lower pH values (Scheme 29.20). A similar mechanism including the two-electron reduction of the azo group to the hydrazo group in basic aqueous solutions and the four-electron process with N–N bond cleavage yielding substituted anilines in acidic media was suggested for many azo dyes (cf., the review [2]). Because the N–N bond cleavage is catalyzed by acids and bases, one four-electron polarographic wave was observed not only for XVI in acidic media but also in alkaline media for its derivatives with electron-withdrawing substituents [93]. However, the four-electron reduction of N,N-dimethyl-4amino-4′-hydroxyazobenzene in water–methanol mixtures in the whole pH range determined by coulometry was explained [94] by fast disproportionation of the hydrazo derivative formed in the first two-electron exchange yielding two derivatives of aniline and the parent azo compound that is further reduced. A small difference in reduction potentials of metastable cis and stable trans isomers of XVI was observed on mercury [95,96] and pyrolytic graphite electrodes [91]. In DMF at low temperatures, cis-azobenzene is reduced in the one-electron process at potentials 60 mV more negative than for the trans-isomer and this process is followed by rapid isomerization of the radical anion [95]. Both isomers are reduced at the same potential in pure aqueous solutions (where they are strongly

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1140

Organic Electrochemistry

adsorbed) and in pure ethanol due to fast isomerization [96]. The same potential for both isomers was also found in pure ethanol up to pH 11 (where there is no adsorption). However, in alkaline ethanol solutions with pH > 11 cis-azobenzene is reduced at more positive potentials [96], which was confirmed later using a channel electrode [97]. More recently, photomodulated voltammetry was used to determine accurate reduction potentials for in situ generated cis-azobenzene in eight aprotic solvents [98]. Solvent effect on potentials for both isomers was described by the empirical solvent parameter π*. The proposed mechanism involved reversible heterogeneous and solution electron transfers for both isomers as well as isomerization equilibria [98]. Substituent effects on the electrode behavior of substituted azobenzenes are mainly caused by affecting its basicity. Namely, electron-withdrawing substituents (e.g., –COOH, –SO3H, –CN) lowering the basicity of the azo group facilitate its two-electron reduction, whereas electron-donating substituents (e.g., –OH, –NH2, –N(Me)2) facilitate protonation and cleavage of the N–N bond [92]. Substituent effects on the reduction potentials of azobenzenes and related compounds in different media were intensively investigated and described in terms of the Hammett equation or related approaches. For example, polarographic data were used to calculate substituent constants for a series of pyridylazo and thiazylazo dyes in 50% ethanol [99]. Some other examples are reviewed in Reference 4. On the other hand, the presence of the o-nitro substituent leads to cyclization. In particular, the mechanism of preparative formation of substituted benztriazole-N-oxide and next benztriazole by the electroreduction of o-nitrophenylazo-p-cresol in basic solutions was discussed by Bourgeois et al. [100]. The electroreduction of many azobenzene derivatives with a number of substituents as well as some other aromoatic azo compounds was widely studied in aprotic solvents and some interesting results will be mentioned further on. Reduction of azobenzene-4-sulfonic acid in MeCN involves autoprotonation and thus four separate consecutive waves were observed for azonium species and sulfonate ions formed in a sequence of protonation equilibria [101]. Two one-electron reversible steps were observed for the reduction of 2,2′,4,4′,6,6′-hexanitroazobenzene in MeCN and hexanitrohydrazobenzene was the final product in the presence of proton donors [102]. The electroreduction of substituted 2-nitroazobenzene and 2-nitrohydrazobenzene in DMF with n-Bu4NClO4 yielding 1,3-dihydroxy-2-phenylbenztriazole was reported [103]. However, in the presence of sodium and lithium cations, 2-phenyl-benztriazole-N-oxide (which could be further reduced to 2-aminoazobenzene) was the two-electron reduction product. The cathodic behavior of a series of aromatic azo compounds, including azonaphthalenes and azoanthracenes, is similar to that of XVI. In particular, it was well established that the first one-electron step yielding radical anions is reversible [87]; for example, it was observed for 1,1′-azonaphthalene and 4,4′-azopyridine in DMF [82], 2,2′- and 4,4′-azopyridine, 2-, 3- and 4-phenylazopyridine in MeCN [86]. The formation of ion pairs between radical anions of 2,2′-azonaphthalene with Li+ cations in DMF and kinetics of their disproportionation were described [104]. For a series of 2-arylazaanthraquinones in benzonitrile solutions, the reduction of the azo group to the hydrazo group was reported in the first step before the reduction of the quinone moiety [105]. The reaction entropy of the formation of radical anions of aromatic azo compounds in MeCN was measured and related to the reactant size and the spin density at the nitrogen atom of a radical anion [106]. On the other hand, the reduction of 1-azuleneazobenzene and other azo-1-azulene compounds in MeCN exhibited [107] a quasi-reversible first electron transfer (which produces unstable radical anions generating polymeric films) and an irreversible second step. Finally, two respective amines were formed by N–N bond cleavage, but unstable 1-aminoazulene could only be produced at very low temperature [107]. Thanks to a good understanding of the behavior of the azobenzene–hydrazobenzene redox couple, it was possible to broaden the scope of its use and look for some new applications. The electrochemical study on Au(111) of self-assembled monolayers containing azobenzene [108] can be mentioned here as well as a recently proposed azobenzene-functionalized electrode [109], which

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1141

Aromatic Nitrogen–Containing Compounds

O

O

–

O

N Ph

(LXIX)

OH

+e

+e Ph N

–

Ph N

N Ph

Ph N N Ph

(LXIX– )

SH, –S–

Ph N N Ph –OH–

(LXIX2–) XVI2–

+e

XVI–

+e

Ph N

N Ph

(XVI)

SCHEME 29.21

Mechanism for the electroreduction of azoxybenzene LXIX in DMF.

could act either as a cathodic molecular rectifier stopping the anodic current or as an anodic rectifier. The CV behavior of azobenzene microcrystals attached to a graphite electrode can also be mentioned [110]. Azobenzene itself could not be oxidized in acetate buffer solutions at a pyrolytic graphite electrode [111]. However, after substitution by electron-donating groups, like the N,N-dimethyl group, the oxidation of the substituent was observed at graphite [111] and platinum disk [112] electrodes, without affecting the azo group. For azo-azulene compounds [107], the irreversible anodic oxidation of the azulene moiety having a higher electron density occurs more easily than that of the azo group. On the other hand, it was recently proven that the oxidation of azobenzene derivatives in aprotic solvents occurs easily. For example, 3-nitroazobenzene gives two one-electron irreversible CV peaks in MeCN and the first electron transfer from 4-nitro-4′-methoxy-azobenzene in dry methylene chloride with the formation of radical cation is reversible [113]. The electroreduction of azoxybenzene (LXIX) in DMF starts from the one-electron step, resulting in the reversible formation of the respective radical anions [114] (Scheme 29.21). Similar behavior was also found for its derivatives, and the formation of radical anions of 2,2′-dimethylazoxybenzene was recently confirmed by in situ EPR experiments [115]. The second electron transfer (also reversible according to recent investigations [115]) is followed by a rapid reaction of dianions with the final formation of azobenzene XVI, which is immediately reduced because its formal potential corresponds to more positive values, before the first reduction peak of LXIX [114,115]. Thus, the azobenzene peak is observed in CV curves during the second scan [115]. However, after the addition of an effective proton donor, a “pre-peak” appears also in the first scan. Careful analysis of CV behavior in dry DMF and after the addition of proton donors has allowed Simonet and coworkers to identify [115] the follow-up reaction as the protonation of dianions by acidic impurities or solvent molecules, SH. This reaction is followed by fast elimination of OH− ions [115]. The general mechanism shown in Scheme 29.21 can explain different CV behavior reported under different experimental conditions (using various scan rates and solvents of varying level of impurities): the reduction of intermediate radical anions XVI− • was observed by Lipsztajn et  al. [114] as a separate third peak, whereas Simonet et al. [115] observed for substituted azoxybenzenes one irreversible three-electron peak including the reduction of a radical anion to a dianion as well as the two-electron reduction of substituted azobenzene. However, in the presence of alkali metal cations of the supporting electrolyte, the CV behavior was seriously changed due to the follow-up reactions of ion associates [114]. On the other hand, in protic solvents azoxybenzene and its derivatives undergo a four-electron and four-proton reduction to hydrazobenzenes [91,92,116]. This process is more difficult than the analogous reduction of azobenzenes that can be observed at more positive potentials in the second scan. Linear dependencies of half-wave potentials on pH for the four-electron reduction of LXIX in acidic and alkaline media were determined [91]. In acidic solutions with pH < 4, the polarographic reduction of azoxybenzene exhibits a second two-electron wave [91,92] that corresponds to the reduction of diprotonated hydrazobenzene cations to aniline in the same process as is shown

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1142

Organic Electrochemistry

in Scheme 29.20. This process is difficult to observe for the reduction of original hydrazobenzene VIII because of its fast rearrangement to benzidine [92]. Moreover, VIII is not reducible at the dropping mercury electrode in less acidic media. During the reduction of LXIX in MeOH with LiCl as the supporting electrolyte, only one four-electron wave was observed [117]. The detailed reduction mechanism of substituted azoxybenzenes at the controlled potential mercury pool cathode was investigated by Hazard and Tallec [116]. In particular, it was shown that p-, m-, and o- nitro and dinitro derivatives undergo reduction in buffered solutions forming in the first four-electron process the respective hydrazo compounds. Complicated schemes for further reactions including the reduction of nitro groups and the disproportionation of ortho- and para-substituted intermediates to substituted anilines as well as the formation of 2-phenyl-benztriazoles from o-nitro and o,o′-dinitro compounds were proposed [116]. Quantitative reduction of azoxytoluene to hydrazotoluene on a stainless steel electrode in alkaline ethanol–water solutions [118] can be pointed out as an example of the use of azoxy compounds in preparative electrolysis.

IV. ARyL DIAZONIUM SALTS Cathodic reduction of aryl diazonium salts in protic and aprotic media has been studied very intensively for many years because of potential applications of such processes in preparative electrochemistry like phenylation of various compounds and more recently in the derivatization of the electrode surface through the electrografting of aryl groups. The last process is discussed in detail in Chapter 42 and will not be discussed here. However, the general mechanism of the formation of aryl radicals from diazonium salts at electrodes and their further reactivity will be briefly reviewed. In early polarographic studies of benzenediazonium salts (also bearing some substituents), it was established [119–123] that the reduction process in aqueous acidic solutions exhibits two waves. The first one is pH independent and corresponds to a one-electron process as confirmed by coulometric data [119–122]. The last observation as well as the isolation of diaryl mercury and aryl-mercuric chlorides [119,121] indicated the formation of radicals as intermediates. It was suggested [121,122] that these intermediates are diazenyl radicals PhN2•, which adsorb on a mercury electrode and decompose to phenyl radicals in further reactions. The second three-electron and pH dependent process, complicated by side reactions, affords in acidic solutions phenylhydrazine PhNH–NH2 as the final product during preparative electrolysis [122,123]. In aprotic solvents like sulfolane [124] and acetonitrile [125–128], the first electroreduction step of benzenediazonium tetrafluoroborates (LXX) is also a one-electron process resulting in the formation of phenyl radicals. This was confirmed indirectly by identification of final products. For example, the isomer ratio of biaryls formed during the electrolysis in the presence of a number of aromatic compounds and partial rate factors of these reactions were in accordance with those obtained for well-known free radical phenylation [125]. Moreover, the low-temperature reduction of LXXa in MeCN at a mercury pool cathode in the presence of α-phenyl-N-tert-butylnitrone showed a strong ESR signal of the adduct with the phenyl radical [126]. However, the slope of the polarographic wave of the reduction of LXXb in sulfolane suggested a reversible process [124]. Thus, the formation of intermediate PhN2• radicals in the first step followed by the C–N bond cleavage in the second step was assumed in earlier papers [124,125]. However, when using glassy carbon electrodes and small concentration (below 1 mM) of LXX in order to diminish fouling of the electrode surface, it was recently established [109] that the first step is a concerted electron transfer and bond cleavage (see also Chapters 13 and 14). This mechanism, shown in Scheme 29.22, was evident from the observations that are presented further on. The reduction of LXXc in MeCN exhibits an irreversible broad CV peak, followed by the reversible cathodic and anodic peaks of nitrobenzene [127], which were formed after the C–N bond cleavage. The width of the first peak equal to 180 mV is characteristic of concerted processes. Similar peak widths were found [128] for LXXa and LXXb

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1143

Aromatic Nitrogen–Containing Compounds +e X +

X

N N BF4–

+e –BF4–

N2 + X

C

C +H X

(LXXa) X = H (LXXb) X = Me (LXXc) X = NO2

SCHEME 29.22 Pathways for the electroreduction of benzenediazonium tetrafluoroborates LXX in aprotic media.

giving the apparent transfer coefficients α = 0.34 and 0.30, respectively. These low values of α indicate a concerted mechanism with great reaction driving force that is necessary to overcome very strong activation free energy, which includes additional contribution from the bond energy [129]. Moreover, a very small second peak was observed at more negative potentials in MeCN and DMSO [128]. It corresponds to the reduction of phenyl (or substituted phenyl) radicals to anions. The simulation of CV curves permitted the authors to confirm the proposed mechanism and to determine the standard potential of the Ph·/Ph− couple in MeCN as being equal to 0.05 V versus SCE. Moreover, Ph· radicals undergo a competitive reaction of a fast hydrogen atom transfer from solvent molecules diminishing the second peak (Scheme 29.22). A good linear correlation of peak potentials for a series of aryldiazonium salts against the Hammett substituent constants supported the validity of the concerted mechanism for other compounds [128]. Application of phenyl radicals, generated electrochemically by the reduction of LXXa in MeCN, to phenylation of naphthalene and benzene and its derivatives with total yields of different isomers up to 33% was reported [125]. The low yields were caused by the competitive H atom abstraction from the solvent by phenyl radicals. On the other hand, the cyclization of free radicals generated by the reduction of diazonium salts of 2-amino-cinnamic acid derivatives at a mercury electrode in MeCN affords phenanthrene-10-carboxylic acids with 80–96% yields [130]. Some reactions of benzenediazonium cations formed in MeCN solutions during the anodic oxidation of triazenes will be considered in Section V.

V. AROMATIC TRIAZENES Aromatic triazenes can be electrochemically reduced or oxidized. Most investigations performed around 1972 were reviewed by Iversen [131] who collected half-wave potentials obtained as a result of the polarographic reduction of 1,3-diphenyltriazene (LXXI) and its phenyl substituted derivatives in aqueous-alcoholic solutions of different pH. After extensive research, Holleck and Kazemifard concluded [132] that aniline and phenylhydrazine (or their derivatives in the case of substituted LXXI) are formed as the final products of the overall four-electron reduction process in neutral 50% aqueous methanol as is shown in Equation 29.5. Only one polarographic wave was observed in methanol solutions containing organic acids, but the same final products were found [133]. A tentative scheme of reduction pathways was proposed [132], but details were not proven. C6H5–N = N–NH–C6H5 + 4e + 4H2O → C6H5–NH2 + C6H5–NH–NH2 + 4OH−

(29.5)

Peak potentials for the one-step irreversible reduction and the two-step irreversible oxidation of a series of asymmetric 1-phenyl-3-alkyltriazenes in MeCN were reported [134]. No signals were found in simultaneous ESR experiments indicating a rapid consecutive reaction of primary radical products. However, using spin-trapping some radical products of the fragmentation started by electrode processes were identified. The obtained results suggest two tautomeric forms of the investigated triazenes [134].

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1144

Organic Electrochemistry Me

+

+

N

N N N

Me –e

(LXXIIb

Me –Other – Products Me

Me

+

N

N N

Me

2+)

(LXXIVa) +

R’ R

N N N

–e

+

R

R’

N N N Me

Me (LXXIIa) R’ = Me, R = H (LXXIIb) R’ = Me, R = N(Me)2 (LXXIIc) R’ = Me, R = OEt (LXXIId) R’ = Me, R = OMe (LXXIIe) R’ = Me, R = Me (LXXIIf) R’ = p-C6H4OMe, R = N(Me)2 (LXXIIg) R’ = p-C6H4OMe, R = NO2 (LXXIIh) R’ = p-C6H4OMe, R = OMe

– N2

Me

(A)

Me

Me N

(LXXII+ )

N Me (LXXIIIb+ )

(B) (D)

(C)

N

OMe + R

+

N N

Me +

+

N

Me +e –e

Me

–

N N

–

R

Me N

+

Me +e –e Me

+

R

N N

(LXXIV)

+

+N Me

NH2

Me-C CH-CN (LXXVI) R

N N C-CN

R

(LXXIVa) R = N(Me)2 (LXXIVb) R = NO2 (LXXIVc) R = OMe

+

N N (LXXIVd) R = OEt (LXXIVe) R = Me (LXXIVa) (LXXIVb) MeO (LXXIVc)

Me N

OMe

N Me (LXXV) –e LXXV+

H2N-C-Me (LXXVII) R = OEt, OMe, Me, H, Br, CF3

SCHEME 29.23 Proposed pathways for the anodic oxidation of aromatic triazenes LXXII in different media.

The anodic oxidation of 1-phenyl-3,3-dimethyltriazene (LXXIIa) in nitromethane solutions indicated [135] the cleavage of the primary radical cation to phenyldiazonium ion as the main reaction route, but this cleavage occurs only as a side reaction in MeCN. On the other hand, the oxidation of 1-(p-dimethylaminophenyl)-3,3-dimethyltriazene (LXXIIb) at a platinum-rotating disk electrode in aqueous media showed [136] two waves independent of pH and corresponding to one-electron steps with the formation of the radical cation LXXII+ • and next the dication LXXII2+ as is shown in the top part of Scheme 29.23. However, the experimental detection of intermediates LXXII+ • and LXXII2+ was possible only in aprotic media, and it was reported by Speiser and coworkers in a series of papers [137–140] elucidating the details of the oxidation mechanism of aryl triazenes in MeCN at Pt electrodes. Different pathways and products of the decay of radical cations LXXII+ • were found depending on the nature of substituents. For the first oxidation step of LXXIIb, simulations of CV peaks confirmed the eC mechanism with the reversible electron transfer (rate constant equal to 0.01 s−1) followed by a slow and irreversible reaction [137]. ESR spectra during the electrolysis of LXXIIb at room temperature indicated [137] the expulsion of nitrogen and the formation

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1145

Aromatic Nitrogen–Containing Compounds

N N N H

–e

LXXI+

+ LXXI+ + H2O

2 LXXI + 2 H+

–1/2O2

(LXXI) 2

N N N H H

2 H2N

N N (LXXVIII)

SCHEME 29.24

Anodic oxidation of 1,3-diphenyltriazene LXXI in MeCN.

of N,N,N′,N′-tetramethyl-p-phenylenediamine radical cation (Wursters Blue, LXXIIIb+ •) according to the pathway (A) in Scheme 29.23. Moreover, the rapid decomposition of dications LXXIIb2+ resulted in the formation of p-dimethylaminobenzenediazonium cations (LXXIVa) as the main product (Scheme 29.23). Benzenediazonium cations with different substituents were identified electrochemically (by their reduction peaks) as decay products of radical cations LXXII+ • formed by the oxidation of other 1-(p-substituted-phenyl)-3,3-dimethyltriazenes. Thus, the existence of the bond cleavage between N(2) and N(3) was confirmed [138]. On the other hand, 2,7-dimethoxy-5,10dimethyl-5,10-dihydrophenazine radical cation (LXXV+ •) was identified using ESR and ENDOR spectroscopy during the electrochemical as well as chemical oxidation of LXXIIf, LXXIIg, and LXXIIh [138] according to the pathway (B) in Scheme 29.23. Finally, evident ESR spectra of radical cations LXXII+ • were obtained using a rapid-scan spectrometer in situ for the anodic oxidation of LXXIIb and LXXIIc as well as for LXXIIf and LXXIIg [139]. Uv–Vis spectra of LXXII+ • were also recorded during the voltammetric experiments. In particular, for the oxidation of LXXIIc, separate absorption bands were observed for the reactant (at 292 and 326 nm), for the radical cation (at 440 nm and a flat maximum above 600 nm), and for the product 4-ethoxybenzenediazonium ion LXXIVd (at 317 nm) [139]. All experimental data support the eCe mechanism with the bond cleavage as the chemical step that can occur either by affording directly LXXIV and dimethylaminyl radical according to the pathway (C) or by the formation of diazenyl radical and its further oxidation through the pathway (D) [139]. Further reactions of diazonium ions LXXIV generated in the anodic oxidation in MeCN solutions include first of all the coupling with 3-aminocrotoninitrile (LXXVI) generated simultaneously at the cathode. This reaction produced enamines LXXVII as the main product with estimated yields of 25–30% [140]. However, for LXXIIg with the electron-withdrawing NO2 substituent, 4-nitrobenzenediazonium ions LXXIVb were formed only after the oxidation to dications LXXII2+ [139] and nitrobenzene is the main product of the decay of LXXIVb [140]. On the other hand, the extensive study [135] of the oxidation of 1,3-diphenyltriazene LXXI in MeCN solutions using voltammetry at a platinum-rotating electrode and the CV technique proved the formation of a radical cation in the one-electron step and its back reduction in the homogeneous reaction with residual water (Scheme 29.24). Further protonation of LXXI and next a rearrangement result in the formation of p-aminoazobenzene (LXXVIII).

ACKNOWLEDgMENT The author wishes to express his indebtedness to the late Stanisław J. Jaworski, PhD student at the University of Gdańsk, for his kind help in the preparation of schemes.

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REFERENCES 1. Jaworski, J. S.; Kalinowski, M. K. Electrochemistry of anilines. In The Chemistry of Anilines, Part 2; Rappoport, Z., ed.; Wiley Interscience: Chichester, U.K., 2007; pp. 871–929. 2. Simonet, J.; Gueguen-Simonet, N. Cathodic and anodic behaviour of organic compounds possessing – N=N– or >N–N< linkages. In PATAI’s Chemistry of Functional Groups; Vol. 2; Patai, S., ed.; Wiley Interscience: Chichester, UK., 1997; pp. 391–463. 3. Reed, R. C.; Wightman, R. M. Aromatic amines. In Encyclopedia of Electrochemistry of the Elements: Organic Section; Vol. XV: Derivatives of Ammonia; Bard, A. J., Lund, H., eds.; Marcel Dekker: New York, 1984; pp. 3–54, 100–116. 4. Stradins, J. P.; Glezer, V. T. Azo, azoxy, and diazo compounds. In Encyclopedia of Electrochemistry of the Elements: Organic Section; Vol. XIII; Bard, A. J., Lund, H., eds.; M. Dekker: New York, 1979; pp. 163–208. 5. Genies, E. M.; Lapkowski, M. J. Electroanal. Chem. 1987, 236, 189–197. 6. Yang, H.; Bard, A. J. J. Electroanal. Chem. 1992, 339, 423–449. 7. Geniès, E. M.; Lapkowski, M.; Penneau, J. F. J. Electroanal. Chem. 1988, 249, 97–107. 8. Mu, S.; Kan, J. Electrochim. Acta 1996, 41, 1593–1599. 9. Sharma, L. R.; Manchanda, A. K.; Singh, G.;Verma, R. S. Electrochim. Acta 1982, 27, 223–233. 10. Desideri, P. G.; Lepri, L.; Heimler, D. J. Electroanal. Chem. 1971, 32, 225–234. 11. Male, R.; Allendoerfer, R. D. J. Phys. Chem. 1988, 92, 6237–6240. 12. Petr, A.; Dunsch, L. J. Phys. Chem. 1996, 100, 4867–4872. 13. Zimmermann, A.; Künzelmann, U.; Dunsch, L. Synth. Metals 1998, 93, 17–25. 14. Stassen, I.; Hambitzer, G. J. Electroanal. Chem. 1997, 440, 219–228. 15. Cases, F.; Huerta, F.; Garcés, P.; Morallόn, E.; Vázques, J. L. J. Electroanal. Chem. 2001, 501, 186–192. 16. Deng, H.; Van Berkel, G. J. Anal. Chem. 1999, 71, 4284–4293. 17. Ohsaka, T.; Okajima, T.; Oyama, N. J. Electroanal. Chem. 1986, 200, 159–178. 18. Oyama, N.; Ohsaka, T.; Shimizu, T. Anal. Chem. 1985, 57, 1526–1532. 19. Hambitzer, G.; Heitbaum, J.; Stassen, I. J. Electroanal. Chem. 1998, 447, 117–124. 20. Yang, H.; Wipf, D. O.; Bard, A. J. J. Electroanal. Chem. 1992, 331, 913–924. 21. Oyama, M.; Higuchi, T. J. Electrochem. Soc. 2002, 149, E12–E17. 22. Rees, N. V.; Klymenko, O. V.; Compton, R. G.; Oyama, M. J. Electroanal. Chem., 2002, 531, 33–42. 23. Yang, H.; Bard, A. J. J. Electroanal. Chem. 1991, 306, 87–109. 24. Andrieux, C. P.; Gallardo, I.; Junca, M. J. Electroanal. Chem. 1993, 354, 231–241. 25. Park, H.; Oyama, M. J. Chem. Soc. Perkin Trans. 2, 2002, 1335–1339. 26. Yoshida, Y.; Fueno, T. J. Org. Chem. 1972, 37, 4145–4147. 27. Pekmez, N.; Pekmez, K.; Yildiz, A. J. Electroanal. Chem. 1993, 348, 389–398. 28. Serve, D. Electrochim. Acta 1976, 21, 1171–1181. 29. Larumbe, D.; Gallardo, I.; Andrieux, C. P. J. Electroanal. Chem. 1991, 304, 241–247. 30. Oyama, M.; Nozaki, K.; Okazaki, S. Anal. Chem. 1991, 63, 1387–1392. 31. Steckhan, E. Top. Curr. Chem. 1987, 142, 1–69. 32. Bacon, J.; Adams, R. N. J. Am. Chem. Soc. 1968, 90, 6596–6599. 33. Desideri, P. G.; Lepri, L.; Heimler, D. J. Electroanal. Chem. 1974, 52, 105–114. 34. Simon, P.; Farsang, G.; Amatore, C. J. Electroanal. Chem. 1997, 435, 165–171. 35. Amatore, C.; Farsang, G.; Maisonhaute, E.; Simon, P. J. Electroanal. Chem. 1999, 462, 55–62. 36. Kádár, M.; Nagy, Z.; Karancsi, T.; Farsang, G. Electrochim. Acta 2001, 46, 1297–1306. 37. Kádár, M.; Nagy, Z.; Karancsi, T.; Farsang, G. Electrochim. Acta 2001, 46, 3405–3414. 38. Goto, M.; Otsuka, K.; Chen, X.; Tao, Y.; Oyama, M. J. Phys. Chem. A 2004, 108, 3980–3986. 39. Streeter, I.; Wain, A. J.; Thompson, M.; Compton, R. G. J. Phys. Chem. B 2005, 109, 12636–12649. 40. Wawzonek, S.; McIntyre, T. W. J. Electrochem. Soc. 1967, 114, 1025–1029. 41. Hawley, D.; Adams, R. N. J. Electroanal. Chem. 1965, 10, 376–386. 42. Piette, L. H.; Ludwig, P.; Adams, R. N. Anal. Chem. 1962, 34, 916–921. 43. Kapturkiewicz, A.; Jaenicke, W. J. Chem. Soc. Faraday Trans. I 1987, 83, 2727–2734. 44. Fawcett, W. R.; Foss, Jr., C. A. J. Electroanal. Chem. 1989, 270, 103–118. 45. Fernández, H.; Zόn, M. A. J. Electroanal. Chem. 1992, 332, 237–255. 46. Moressi, M. B.; Zόn, M. A.; Fernández, H. Electrochim. Acta 2000, 45, 1669–1682. 47. Hand, R. L.; Nelson, R. F. J. Electrochem. Soc. 1978, 125, 1059–1069. 48. Jackowska, K.; Bukowska, J.; Kudelski, A. J. Electroanal. Chem. 1993, 350, 177–187. 49. Widera, J.; Grochala, W.; Jackowska, K.; Bukowska, J. Synth. Metals 1997, 89, 29–37.

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Aromatic Nitrogen–Containing Compounds 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. 61. 62. 63. 64. 65. 66. 67. 68. 69. 70. 71. 72. 73. 74. 75. 76. 77. 78. 79. 80. 81. 82. 83. 84. 85. 86. 87. 88. 89. 90. 91. 92. 93. 94. 95. 96. 97. 98. 99. 100. 101. 102. 103. 104.

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30

Reduction of Nitro Compounds and Related Substrates Ole Hammerich

CONTENTS I. Introduction ........................................................................................................................... 1150 II. Nitro Compounds ...................................................................................................................1151 A. Aromatic Nitro Compounds...........................................................................................1151 1. Formation and Properties of the Radical Anions and Dianions............................1151 2. Mononitrobenzenes .............................................................................................. 1152 3. Other Mononitroarenes......................................................................................... 1166 4. Di- and Trinitroarenes .......................................................................................... 1166 5. Hetereoaromatic and Related Mono- and Dinitro Compounds.............................1170 B. Aliphatic Nitro Compounds ...........................................................................................1170 1. Formation and Properties of the Radical Anions ..................................................1171 2. Nitroalkenes and Related .......................................................................................1172 3. Mononitroalkanes and Related ..............................................................................1174 4. Dinitroalkanes .......................................................................................................1176 5. Trinitroalkanes ......................................................................................................1178 6. Tetranitroalkanes ...................................................................................................1179 7. α-Halonitroalkanes ................................................................................................1179 8. β-Substituted Nitroalkanes ................................................................................... 1180 9. α,β-Disubstituted Nitroalkanes..............................................................................1181 III. Nitroso Compounds ...............................................................................................................1181 A. Aromatic Nitroso Compounds .......................................................................................1181 1. Formation of and Properties of the Radical Anions and Dianions .......................1181 2. Routes to Phenylhydroxylamines and Anilines .....................................................1182 3. Routes to Azoxybenzenes, Azobenzenes, Hydrazobenzenes, and Benzidines .....1183 4. Other Reactions of Nitrosobenzenes .................................................................... 1184 5. Nitronitrosobenzenes ............................................................................................ 1184 6. Heteroaromatic and Related Nitroso Compounds .................................................1185 B. Aliphatic Nitroso Compounds .......................................................................................1185 1. Primary and Secondary Nitrosoalkanes................................................................1185 2. Tertiary Nitrosoalkanes .........................................................................................1185 3. α-Halonitrosoalkanes.............................................................................................1185 4. Nitronitrosoalkanes (Pseudonitroles) ................................................................... 1186 IV. Hydroxylamines .................................................................................................................... 1186 A. Aromatic Hydroxylamines ............................................................................................ 1186 B. Aliphatic Hydroxylamines ............................................................................................ 1186

1149

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Organic Electrochemistry

V. N-Nitramines, N-Nitrosamines, N-Nitramides, and N-Nitrosamides ................................... 1187 A. N-Nitramines and N-Nitrosamines ............................................................................... 1187 B. N-Nitramides and N-Nitrosamides................................................................................ 1188 VI. Nitric Acid Esters ...................................................................................................................1189 References .................................................................................................................................... 1190

I. INTRODUCTION The electrochemical reduction of nitro compounds played an important role in the development of organic electrochemistry in the early period around 1900 [1]. For example, it was studies of the stepwise reduction of nitrobenzene that originally led Haber [2] to realize the importance of the potential of the working electrode in electrochemistry and to establish the redox relationship between nitrobenzene, nitrosobenzene, phenylhydroxylamine, aniline, p-hydroxyaniline, azoxybenzene, azobenzene, hydrazobenzene, and benzidine [2,3]. The essential parts are reproduced in Scheme 30.1. This early knowledge of the products that may result from reduction of nitro compounds stems from reactions carried out in aqueous solutions or suspensions and the first insight into the reaction details have resulted from early polarographic investigations, also in aqueous solution. Later this was supplied by studies in nonaqueous and aprotic media. The latter in particular have allowed for the observation of the primary electrode products, the radical anions and dianions. In this chapter, we will discuss the reduction of organic nitro compounds and some of the reduction products. As seen in Scheme 30.1, nitroso compounds are intermediates in the reduction of nitro compounds and their electrochemistry is therefore included under nitro compounds when needed. Electrochemical processes in which the nitroso compounds are the starting materials are treated separately. The electrochemistry of azoxybenzene, azobenzene, hydrazobenzene, and benzidine is only included to the extent necessary for the discussion. A more detailed account of the electrochemistry of these dimers is given in Chapter 29. The readers of this chapter may also find discussions of interest in Chapter 44 and in other electrochemistry books [4–8]. Some preparative aspects have been briefly reviewed [9]. Ph +2e–,+2H+

Ph

NO2 –H2O NO Ph

+2e–,+2H+

Ph

–H2O NHOH

NH2

Ph

Ph

–

O

+2e–,+2H+

+2e–,+2H+ –H2O HO

N+ N

Ph

N

–H2O N

Ph

+2e–,+2H+

NH2

Acidic

Ph

NH

NH

Ph

Conditions H2N

SCHEME 30.1

NH2

Overview of the products that may result from the reduction of nitrobenzene.

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Reduction of Nitro Compounds and Related Substrates

II.

1151

NITRO COMPOUNDS

A. AROMATIC NITRO COMPOUNDS 1. Formation and Properties of the Radical Anions and Dianions The electrochemical reduction of nitrobenzene to a persistent radical anion may be observed in almost any aprotic solvent. Examples include MeCN [10–13], DMF [11,14–20], DMSO [11], CH 2Cl2 [21], THF [11], and propylene carbonate [22] and specialty solvents such as liquid [23] and supercritical ammonia [24] and room temperature ionic liquids [25]. The same is true for many substituted nitrobenzenes [10,11,13,16,19,22,26–57] and other aromatic [19,20,26,58–60] or heteroaromatic nitro compounds [61–68] as long as the structures do not include, for instance, acidic hydrogens that may protonate the radical anions (self-protonation; see Section II.A.2.a.i). Other exceptions include radical anions that owing to a high local spin density undergo (reversible) dimerization as observed, for example, for 1,3,5-trinitrobenzene [19,69,70], 3,5-nitropyridine [70], and 9-nitroanthracene [19,71] and radical anions that carry a potential leaving group such as a halide ion [72–74]. The formation and reoxidation of the radical anions of nitrobenzene and substituted nitrobenzenes in aqueous solution or in mixtures of an organic solvent and water require nonacidic conditions in order to avoid protonation of the radical anion [75–78]. Radical anions have been observed also in aqueous solutions containing micellar components [79]. Further reduction of the nitrobenzene radical anion to a persistent dianion requires the strict absence of proton donors, including residual water [14], or low temperature as in liquid ammonia [23]. Less reactive dianions are observed for nitrobenzenes carrying phenyl substituents [31], one or more extra nitro groups [19,22,35,36,38–47,56,80], or another electronwithdrawing group such as nitroso [48], azo [49], carbonyl [50,51], methylsulfonyl [55], N,Ndialkylsulfamoyl [52], and 3-thioxo-3H-1,2-dithiol-5-yl [81] or a combination of some of these structural features [53]. The dianions of 1,3-dinitrobenzene [40,42], 1,3,5-trinitrobenzene [69], methyl 3,5- dinitrobenzoate [42], and 2,2′,4,4′-tetranitrobiphenyl [46] have been reported to have a biradical nature. Radical trianions and even tetraanions have been observed for tetranitro compounds such as 2,2′,4,4′-tetranitrobiphenyl [46] and 2,2′,4,4′-tetranitrodiphenylmethane [19]. Reduction potentials have been recorded for a number of substituted nitrobenzenes [10,13,82,83], and not unexpectedly electron-withdrawing substituents make the nitro compounds easier to reduce and electron-donating substituents make them more difficult to reduce. Reduction potentials have also been recorded in DMF for a large series of nitroarenes [19]. The experimental values of E°′ have been found to correlate linearly with calculated values of E LUMO [10,84] illustrating the well-known fact that the changes in solvation energies and other properties associated with electron transfer usually vary gradually along a series of related compounds. Similarly, the relationship between experimental values of E°′ and calculated electron affinities has been discussed [85]. Advances in computer technology and the development of user-friendly software have now made it possible also to obtain redox potentials for the reduction of nitrobenzenes by calculation, typically at the DFT level of theory [86,87] (see Chapter 6 for details). The solvation energy changes associated with the reduction of aromatic nitro compounds to the radical anions in aprotic solvents are in the range −40 to −70 kcal mol−1 [11]. The entropy of formation of the radical anions in MeCN has been determined for a series of mono-, di-, and trinitrobenzenes and alkyl substituted derivatives and was found to be influenced by both steric and electronic effects. Values are typically in the range −10 to −20 cal K−1 mol−1 [35]. The thermodynamics for the successive electron transfers for a series of polynitro compounds addressing the effects of internal properties and the effects of solvation and ion-pairs have been investigated as well [36,45]. Of course, any interaction that stabilizes the radical anion and dianion in solution relative to the precursor will make reduction easier and not only stabilization

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1152

Organic Electrochemistry

by solvation, but also by ion-pairing [22,39,45,60,79,88–97] and hydrogen-bonding [21,38,98], facilitates reduction. In general, ion-pairing is stronger with alkali metal ions than with tetraalkylammonium ions, dications are more effective than monocations, and the dianions of the nitro compounds are more affected than the radical anions. For example, the disodium salt of the nitrobenzene dianion has been isolated from MeCN/NaClO 4 solutions [99] and ion pairs with Mg2+, Ca 2+, and Ba 2+ have been observed as insoluble deposits at the working electrode surface [89]. The first electron transfer resulting in the formation of the radical anion is reversible during slow sweep cyclic voltammetry. Impedance measurements at mercury cathodes in MeCN have shown [30,100–102] that the values of the heterogeneous electron transfer rate constants, k s, are larger than approximately 2  cm s−1 unless substitution in the 2,6-positions hinders coplanarity of the nitro group and the aromatic ring. In that case, the values of k s are lower, typically in the range 0.1 < k s/cm s−1 < 0.8 illustrating the importance of the Marcus inner reorganization energy in determining the kinetics of the electron transfer process [100]. Smaller values than this have been observed for sulfur containing nitro compounds; this has been attributed to adsorption effects [103]. The Marcus outer reorganization energy, related to the solvation energy changes, appears to be a significant factor in determining k s in ionic liquids [30], where also diffusion coefficients are observed to be much smaller than in, for instance, MeCN [29]. The reversibility of the first electron transfer process makes the radical anions suitable as electron donors in redox catalysis [104,105]. We are not aware of quantitative studies of the second electron transfer for mononitrobenzenes. For dinitrobenzenes, the values of k s are typically around 0.01 cm s−1 [80]. UV/Vis [16–18,20,26,31,32,59,69], IR [38], and ESR [12,13,27,33,34,42,44,57,58,61,62,64,65,74, 77,79,106–109] spectra have been recorded for the radical anions of a number of aromatic and heteroaromatic nitro compounds, and from the ESR spectra of the planar systems, it is evident that the spin is well distributed over the entire molecule. More localized spins are observed for compounds in which substituents in the 2- or 2,6-positions force the nitro group out of the plane defined by the aromatic ring. The acidity in DMSO of a series of nitrobenzene radical anions, such as nitrophenol radical anions, has been determined via a thermochemical cycle and the pKa values were found to be typically 5–10 units higher than those for the neutral substrates [37]. 2. Mononitrobenzenes The electrochemistry of compounds in which a single nitro group is attached to a benzene ring or a substituted benzene ring, including nitrophenyl substituted heterocyclic compounds, is discussed in this section. Mechanism studies, mostly focused on protonations, are presented first and after this some preparative aspects of nitrobenzene reduction are discussed. a. Routes to Nitrosobenzenes, Phenylhydroxylamines, and Anilines i. Mechanisms of Protonation in Aprotic Solvents As it appears from Section II.A.1, the acidity of the residual water present in the commonly used aprotic solvents is not sufficient to protonate the nitrobenzene radical anion at the time scale of slow sweep voltammetry. This is not so with the more basic dianion. The voltammetric response of nitrobenzene in DMF [17,57,110,111] propylene carbonate [22] and liquid ammonia containing 2-propanol [23] is composed of a one-electron process leading to the radical anion (Equation 30.1), and, at lower potentials, a three-electron process leading to phenylhydroxylamine (Equation 30.2). The anion of the latter, Ph-NHO−, was detected by its oxidation peak during CV [23]. +e– Ph

© 2016 by Taylor & Francis Group, LLC

NO2

–e–

Ph

NO2 –

(30.1)

1153

Reduction of Nitro Compounds and Related Substrates

Ph

Ph

+e–

NO2 –

NO

–e–

Ph

+e–

–

–e–

Ph

NO2– 2

NO2–

OH

H2O Ph

–OH–

Ph

+e– –OH–

O–

H2O –OH–

N

Ph

NO

–e–

(30.2) H2O

NHO–

–OH–

Ph

NHOH

Nitrosobenzene is more easily reduced than nitrobenzene and the nitrosobenzene dianion is therefore formed directly at the potential required to produce the nitrobenzene dianion. However, the reaction scheme may be slightly more complex than shown here. For instance, solution electron transfers, such as the reduction of nitrosobenzene by the nitrobenzene radical anion, are not included and it is also known that nitrosobenzene and phenylhydroxylamine under strongly basic conditions exist in equilibrium with the nitrosobenzene radical anion [112,113]. See Equation 30.13. The unknown concentrations of residual water in the commonly used aprotic solvents make it almost impossible to control the acid–base reactions that accompany the electrochemical reduction of aromatic nitro compounds in these solvents. For that reason, and for gaining more insight into the protonation of the radical anions, a number of studies have addressed the effect of addition of acids stronger than water [22,23,28,111,114–117]. The general observation is that a pre-wave develops at the expense of the original one-electron wave with increasing concentrations of acid and grows to the height of a 4F process. A variety of organic acids such as p-toluenesulfonic, trichloroacetic, o-phthalic, salicylic and benzoic acid have been used and the results indicated that the reduction proceeds via a hydrogen-bonded complex formed between the nitrobenzene and the acid in the double layer [114–116]. Hydrogen-bonded complexes have been reported also for the radical anions [21,92,118]. Protonation leads to the hydroxylamine [111,114], here illustrated by the reduction of p-chloronitrobenzene (Equation 30.3) [114]. This is similar to what is observed in aqueous solution (see Section II.A.2.a.ii). The sequence of the electron and proton transfers is usually not addressed. +4e–,+4H+ Cl

NO2

–H2O

Cl

NHOH

(30.3)

A variant of the protonation reaction is observed when the nitrobenzene itself carries an acidic function such as a phenol or a carboxylic acid group. In those cases, and in the absence of deliberately added acids, the primarily formed radical anion is protonated by the substrate in a socalled self-protonation or father–son reaction, here illustrated by the reduction of m-nitrophenol to m-hydroxylaminophenol in DMSO (Scheme 30.2) [33,119]. This particular reaction sequence leads to an overall 0.8F stoichiometry (Equation 30.4). 5 m-HOC6H4NO2

+4e–

–

m-HOC6H4NHOH + 4 m- OC6H4NO2 + H2O

(30.4)

A similar reaction scheme describes the reduction of nitrobenzoic acids [120,121], of nitroimidazoles and nitroindoles [63,122] and of nitrobenzenesulfonamides [123]. An additional complexity is observed for p-nitrophenol [25,34,119,124]. In this case, the p-hydroxylaminophenol suffers rapid dehydration to a quinone imine that under the conditions is further reduced to p-aminophenol requiring an additional two electrons and two protons (see Section II.A.2.c.iv). Accordingly, the stoichiometry in this case, and for o-nitrophenol, changes to an overall ∼0.85F process (Equation 30.5). 7 p-HOC6H4NO2

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+6e–

–

p-HOC6H4NH2 + 6 p- OC6H4NO2 + 2H2O

(30.5)

1154

Organic Electrochemistry +e– m-HOC6H4NO2

–e–

m-HOC6H4NO2 – –

m-HOC6H4NO2 – + m-HOC6H4NO2

m-HOC6H4NO2H + m- OC6H4NO2

m-HOC6H4NO2H + m-HOC6H4NO2 – m-HOC6H4NO2H– + m-HOC6H4NO2

m-HOC6H4NO + m-HOC6H4NO2 – m-HOC6H4NO

–

+ m-HOC6H4NO2

m-HOC6H4NO2H– + m-HOC6H4NO2 –

m-HOC6H4NO + m- OC6H4NO2

–H2O

m-HOC6H4NO

–

+ m-HOC6H4NO2 –

m-HOC6H4NOH + m- OC6H4NO2

m-HOC6H4NOH + m-HOC6H4NO2 –

m-HOC6H4NOH– + m-HOC6H4NO2

m-HOC6H4NOH– + m-HOC6H4NO2

m-HOC6H4NHOH + m- OC6H4NO2

–

SCHEME 30.2 Self-protonation mechanism for the reduction of m-nitrophenol in DMSO.

In contrast, the radical anions of 3-nitro-1,2,4-triazole-5-ones that also carry labile hydrogen atoms have been reported to undergo tautomerism in MeCN [61], but apparently not self-protonation. ii. Mechanisms of Protonation in Water and Other Protic Solvents Studies of the reduction of nitrobenzene in aqueous solutions have been in focus since the dawn of organic electrochemistry, and it was observed early that the polarographic response for nitrobenzene and substituted nitrobenzenes was dependent on pH and on the presence of surfactants. The pH dependence reflects the participation of protons in the reduction process, whereas the effect of the surfactants indicates that adsorption of the starting material and/or intermediates and/or products contributes to the detailed description of the reaction mechanism. The extent to which adsorption plays a role depends not only on the nature of the electrode material [75,125] but also on the structure of the nitro compound, its concentration, and solubility. Often it is assumed that adsorption plays only a minor role at substrate concentrations   9 leads to the corresponding azoxybenzenes as the major product (see Section II.A.2.c.vi). The electrode material is usually of minor importance. Advantageous is also to keep the processing time short, for instance by using a flow cell with a porous cathode, and to continuously remove the product that often separates as a crystalline material from the electrolysis solution [191]. iii.

Anilines in General +6e–,+7H+ Ar

NO2

–2H2O

Ar

NH3+

(6F)

(30.18)

In anilines, the nitrogen atom is in its lowest oxidation state, and the direct electrochemical conversion of nitrobenzenes requires a relatively low pH to achieve reduction beyond the hydroxylamine stage and often also the addition of small amounts of zinc, copper, tin, or other salts [193–196]. However, as already pointed out, earlier rearrangement of the hydroxylamine to the p-aminophenol to some extent is unavoidable under the strongly acidic conditions, and therefore mixtures of aniline and p-aminophenol usually result from the direct reduction of nitrobenzenes in acid [195] unless the starting material has substituents in the p-position. For instance, p-nitrobenzoic acid [197–199] and p-nitrophenetole [200] may be reduced almost quantitatively to the corresponding anilines in strong sulfuric acid at elevated temperatures. Rearrangement is less of a problem if the reduction is carried out indirectly, for example, by using a Ti(IV) [201–203] mediator or a Ti/TiO2-cathode [204–206]. Useful conditions for the close-to-quantitative conversion of nitrobenzenes to anilines include a 30–40% sulfuric acid catholyte containing 2–3% Ti(IV) sulfate and a copper cathode [201,207]. The Ti(III) formed serves as the reducing agent and is likely to assist also in the cleavage of the N–O bond in the hydroxylamine. The product separates as the aniline hydrogen sulfate. The mediated reduction may also be carried out as an ex-cell two-phase process [208–210] in which the Ti(III) species, in this case (C5H5)2Ti+, is generated electrochemically in the aqueous solution and then allowed to react with the nitrobenzene dissolved in CH 2Cl2. When nitrobenzenes substituted in the o-position with an ester, carbonate, amide, or carbamate function are reduced in this way, the resulting aniline rearranges in situ to the N-acylated o-aminophenols [209]. In another indirect approach, nitrobenzenes have been reduced to anilines in an emulsion with solutions of copper, iron, tin, or zinc salts at a high current density. The finely divided and oxide-free metal powders resulting from the electrochemical reduction serve to effectively reduce nitrobenzene to aniline [211,212]. Even though reduction of nitrobenzenes under basic conditions usually leads to azoxy derivatives, reduction in basic aqueous methanol can be directed to produce anilines in high yields provided that catalytically active electrodes such as Devarda copper, Raney nickel or copper preoxidized to Cu(OH)2, are used [213–216]. Even if azoxybenzenes are formed, they are further reduced to anilines under the conditions. iv. Aminophenols and Phenylenediamines The reduction of nitrobenzenes carrying a hydroxy or an amino group in the o- or p-position presents a special case. The initially formed phenylhydroxylamines rapidly eliminate water to form easily reducible quinone mono- or diimines, as illustrated in Equation 30.19 for p-substitution (X ═ OH, NH2 or NMe2), resulting in an overall 6F process [151,152,217,218].

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Organic Electrochemistry

The elimination of water takes place in both acid and base; here is only shown the reaction in acid. The values of the rate constant, k, for dehydration are typically in the range 0.4–1.0 s−1 [219–221]. NH2OH+

NHOH +H+

+2e–,H+

k –H2O

–H+ X

NH2

NH

(30.19)

X+

X

X

The reaction has been studied at gold electrodes in neutral and acidic solutions by surfaceenhanced Raman spectroscopy, and for the p- and o-isomers, the formation of the phenylenediamine could be detected in acidic, but not in neutral solution. An oxygen-gold adsorbate stretching mode was detected between 400 and 430 cm−1 at positive potentials implying perpendicular adsorption via the nitro group [222]. Examples of the process include the pilot plant scale reduction of 4-hydroxy-3-nitrobenzoic acid to 3-amino-4-hydroxybenzoic acid at a copper cathode under basic conditions [223] and the reduction of 4-nitroaniline, also at a copper cathode, in dilute hydrochloric acid to 1,4-phenylenediamine in 93% yield [224]. The reduction of N,N-dimethyl-4-nitroaniline to the diamine at low temperature [225] serves to illustrate the case X ═ N(CH3)2. Alkoxy groups in the o- [226,227] or p-position [217,227] suffer partial hydrolysis during reduction of the nitrobenzenes. For instance, reduction of a suspension of o-nitroanisole at an amalgamated monel electrode in diluted sulfuric acid in the presence of CuSO4 leads to a mixture of o-methoxyaniline and o-aminophenol [226] and N,N-dimethyl-4-nitroaniline is converted to 4-aminophenol by reduction at elevated temperature [225]. A p-alkyl substituent may function in the same manner as a hydroxyl group in facilitating loss of water from the phenylhydroxylamine and stabilizing the quinonoid intermediate. Thus, 5-nitroacenaphthene is reduced to 5-aminoacenaphthene in a 6F process in acidic solution [228]. Related to this is the 6F reduction of 2,4,6-triphenylnitrobenzene to 2,4,6-triphenylaniline [31]. p-Aminophenols are the products even when nitrobenzenes without a hydroxy group in the p-position are reduced in strong sulfuric acid at an elevated temperature [229–236]. This is the result of the Bamberger rearrangement that in the “pH region” includes elimination of water from the monoprotonated hydroxylamine to a carbocation that subsequently suffers nucleophilic attack by water (Scheme 30.8) [237]. At higher acidity, in the “Ho region,” also the diprotonated hydroxylamine is involved.

+

OH2+

H

NH

N NHOH

NH

H+

NH2

+H 2 O H

H+ H

N

+

–H+

–H2O OH

NH H

OH2+

+

SCHEME 30.8

Bamberger rearrangement of phenylhydroxylamine to p-aminophenol.

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OH

1163

Reduction of Nitro Compounds and Related Substrates

This synthetically important variant of nitrobenzene reduction was reported already by Gattermann [238] and a variety of conditions have been described in books [7,8] and patents [239]. Good conditions appears to be reduction of an agitated suspension of nitrobenzene in 50% aqueous sulfuric acid at a nickel cathode at 25–60°C after which the product precipitates out in 65% yield [240]. Although successful, a complication is that the formation of p-aminophenol is accompanied by a by-product, 4,4′-diaminodiphenyloxide [241], resulting from a reaction in which the intermediate phenylhydroxylamine attacks itself as a nucleophile in competition with water. The amount of 4,4′-diaminodiphenyloxide increases as expected with increasing concentrations of the phenylhydroxylamine. Unfortunately, the 4,4′-diaminodiphenyloxide is carcinogenic and for this reason p-aminophenol prepared electrochemically is not suitable as a drug intermediate. The reduction of p-nitrotoluene at a mercury cathode in 25% sulfuric acid at 90°C was reported to give the 4F product, 2-methylhydroquinone, apparently resulting from hydrolysis of a Bambergerlike intermediate accompanied by methyl migration [192]. The same type of product was observed for 2-methyl-5-nitrobenzoic acid. v. Other Substituted Anilines Other nucleophiles than water may attack the intermediate carbocation shown in Scheme 30.8. Reduction of nitrobenzene in, for instance, hydrochloric acid leads to a mixture of p- and o-chloroaniline and p-aminophenol [242,243] and reduction of 2-fluoronitrobenzene in HCl/propanol gives 4-chloro-2-fluoroaniline in ~50% current yield [244]. Reduction of p-nitrotoluene in concentrated sulfuric acid results unexpectedly in the formation of 4-amino2′-methyl-5′-nitrodiphenylmethane [238,245]. The origin of this product is not clear. A similar conversion has been observed during the reduction of p-nitroisopropylbenzene in sulfuric acid. In an overall 4F process, 5-amino-3-(4-aminophenyl)-1,1,3-trimethylindan was formed; in this case, loss of water from the initially formed phenylhydroxylamine was suggested to give p-amino-αmethylstyrene that then dimerized to the indan derivative [246]. vi. Azoxybenzenes +6e–,+6H+ 2Ar

NO2

–3H2O

Ar

N+ –N

Ar

(3F)

(30.20)

O–

Azoxybenzenes have been detected as minor products from the reduction of nitrobenzenes under a variety of conditions. The classical preparative conditions include reduction of nitrobenzene under basic conditions [247–252] and preferentially the reactions are carried out such that the product precipitates out and may be collected by filtration after electrolysis. In this way, azoxybenzene escapes further reduction to azo- and/or hydrazobenzene. Examples of successful preparations include the reduction of a well-stirred hot 2.5% NaOH suspension of nitrobenzene at lead or nickel cathodes at low current density. The unreduced nitrobenzene may be removed by steam distillation after which the remaining azoxybenzene solidifies upon cooling. Yields are typically around 85% [249]. Another recipe includes the reduction of a solution of the nitrobenzene in aqueous ethanol containing sodium or ammonium acetate as supporting electrolyte [252,253]. Reduction of 3-nitrobenzophenone at a nickel gauze electrode gives under these conditions the corresponding azoxybenzene in a close-to-quantitative yield [252]. The carbonyl group was not affected. Other examples have been reported in two old reviews [247,248]. Occasionally, azoxybenzenes are formed also in fair yields by reduction of the corresponding nitrobenzene under neutral or acidic conditions. For instance, it was reported that the reduction of 4-nitrodiphenylsulfone and 4-nitrobenzophenone at a mercury cathode at pH 2 under vigorous stirring leads to the azoxybenzenes in 30–40% yield, in addition to the hydroxylamines. The azoxy compounds precipitated out and could be isolated by filtration [157]. Reduction of p-iodosonitrobenzene

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1164

Organic Electrochemistry

at a graphite cathode in neutral 1 M aqueous MgBr2 gives 4,4′-diiodoazoxybenzene in 88% yield. Here it should be noticed that the reduction affects also the substituent that is reduced from the iodoso to the iodo state [254]. An unusual reaction was observed when nitrobenzene was reduced at a mercury cathode in MeCN/Bu4NClO4 saturated with carbon dioxide in the absence of acid. Under these conditions, reduction was reported to give azoxybenzene and bis(tetrabutylammonium) carbonate as the major products (Equation 30.21), accompanied by minor amounts of azobenzene and phenylhydroxylamine [255]. The protons required for the latter were suggested to originate from the Bu4N+ ion since also small amounts of 1-butene was detected. 2

+6e–

NO2 + 3CO2

+

+ 3CO32–

N

N

(30.21)

O–

vii. Azobenzenes +8e–,+8H+ 2 Ar

NO2

–4H2O

Ar

N

N

(30.22)

(4F)

Ar

If, during the reduction of nitrobenzenes in basic solution, the current is allowed to flow until the theoretical amount of charge corresponding to 4F has passed, and preferably at increased temperature to avoid precipitation of the intermediate azoxybenzene, the reduction proceeds to the azobenzenes in yields that are typically around 80–95% [225,248,253,256–258]. Nitrobenzenes may also be reduced to azobenzenes in THF/LiClO 4, and by using magnesium electrodes, the current was alternated at 30 s intervals in order to reduce the loss of magnesium. Yields up to 87% were reported [259]. However, the details of the reaction, including the role of magnesium, are not clear. viii.

Hydrazobenzenes and Benzidines +10e–,+11H+ 2Ar

NO2

–4H2O

Ar

R Benzidine Rearr.

H2N

NH

NH2+

Ar

(5F)

(30.23)

R NH3+

As mentioned earlier, a key issue in the preparation of azoxybenzenes is to avoid further reduction, which may be accomplished by using a two-phase system. Obviously then, when further reduction is wanted there is no reason to conduct the electrolysis in this way; on the contrary and therefore, in order to increase solubility, reductions to hydrazobenzenes and benzidines are often carried out in water/alcohol mixtures or in the presence of solubilizing McKee salts (mixtures of sodium and potassium salts of xylenesulfonates) [260]. Still, reduction in basic aqueous suspensions of nitrobenzene in concentrations up to 5% has been shown to be successful when cathodes having a spongy layer of metals like lead, iron, and zinc were used [261]. Advances have been taken of using a rotating disk cathode made by mild steel with a coating of lead for the purpose [262,263]. In general, the reduction of nitrobenzenes proceeds all the way to hydrazobenzenes when the current is allowed to flow until charge corresponding to 5F has been exchanged [225,248,256,257,260]. Addition of the solutions resulting from preparation of the hydrazobenzenes to an excess of hydrochloric or sulfuric acid causes the immediate rearrangement to the corresponding benzidine

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1165

Reduction of Nitro Compounds and Related Substrates

[257,262,263]. Alternatively, the reduction may be carried out in base, stopped at the azoxy or the azo stage, and then continued under acidic conditions to the benzidine [247,250,251]. In the latter case, care should be taken to avoid the further reduction of the hydrazobenzene to the aniline stage [264]. d. Other Reactions of Mononitrobenzenes i. Cleavages Nitrobenzylhalides undergo C–X cleavage upon electrochemical reduction in MeCN or DMF [72–74]; the resulting nitrobenzyl radicals may dimerize to form 1,2-bis(nitrophenyl) ethanes that are further reduced to the corresponding radical anions and dianions. Reduction of p-nitrophenyl methyl sulfone in DMF produces p-nitrophenolate in a rather complex cleavage reaction [55]. Reductive cleavage of 4-(4′-nitrophenyl)-1,3-dioxolane and 7-nitro-1,3-benzodioxane in basic EtOH/water has been suggested as a key step in a protection-deprotection cycle for ketones, here illustrated by cyclohexanone (Equation 30.24) [78,265]. –

OH HO

–

O

O

O

O

O

(30.24)

O

+e– Products O

NO2

NO2

NO2

O

N +

–

O

Similarly, deprotection by reductive cleavage of the 4-nitrobenzyloxycarbonyl group from the urethane derivatives of primary and secondary amines has been reported [266] (see also Chapter 16). ii. Dimerizations Reversible dimerization of the radical anions derived from aromatic compounds has been observed in a number of cases, as, for example, for 1,3,5-trinitrobenzene [19,69,70], 3,5-dinitropyridine [70], and 9-nitroanthracene [19,71]. Such dimerizations are discussed in Chapter 17. iii. Radical Anions as Nucleophiles Nitroarene radical anions have nucleophilic properties, and these have been put to use in reactions with acetic anhydride [82,267,268] and alkyl halides [269] in MeCN or DMF. The products, N-acetoxy-N-arylacetamides from the reaction with acetic anhydride and N,O-dialkyl-N-arylhydroxylamines from the reaction with alkyl halides, not easy to make otherwise, are formed in 25–85% yields. In both cases, the reactions are believed to proceed via the nitroso compound with the stoichiometry shown in Equation 30.25 for the reaction with acetic anhydride. O Ph

NO2 + 3 (CH3CO)2O

+4e–

CH3 Ph

– CH3 + 4 CH3COO

N

(30.25)

O O

The nitro group in 4-nitrobenzoyl chloride is easier to reduce than the acid chloride part. When the reduction is carried out in MeCOMe/LiClO4, the resulting radical anion is attacked by the acid

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1166

Organic Electrochemistry

chloride and a complex reaction leads to the mixed anhydride of 4-nitrobenzoic acid and 4-(4-nitrobenzoylamino)benzoic acid (Equation 30.26) [270]. NO2

NHCOC6H4NO2-p 1.5F

COCl

(30.26) CO2COC6H4NO2-p

iv. Formation of Heterocyclic Compounds Reduction of nitrobenzenes substituted in the o-position with a side chain that may react with the “reduced nitro group” often leads to ring closure and the formation of heterocyclic compounds. Such conversions are treated in Chapter 34 to which the reader is referred for details. Numerous examples are found also in Torii’s book [6]. Here it suffices to present a few illustrative examples. o-Nitrostyrenes are converted to 1H-indoles by reduction in DMF in the presence of a 10-fold excess of a proton donor such as phenol [271]. Reduction of o-nitroazobenzene [272] and o-nitrohydrazobenzene [273] in DMF leads in both cases to 1,3-dihydroxy-2-phenylbenzotriazole, or in the presence of alkali metal ions, to 2-phenylbenzotriazole-N-oxide. Related to these are the reduction of o-nitrophenylhydrazides to 1,2,4-benzotriazines [274] and of o-nitrophenylazo-p-cresols to benzotriazole-N-oxide [275]. N-(o-Nitrobenzoyl) and N-(o-nitrobenzyl) amides or imides may be reduced to quinazoline derivatives [276,277] and o-nitrobenzoic acid derivatives and o-nitronaphthalimides to isoxazole derivatives [278,279]. In a reduction–oxidation process, o-nitrobenzylamines were converted to the resulting nitroso compounds that subsequently cyclized to indazoles [280,281]. In basic aqueous media, 2-nitroazobenzenes are reduced to 2-arylbenzotriazole-Noxides [282,283]. In DMF, l,3-dihydroxy-2-phenylbenzotriazole is formed from 2-nitroazobenzene [272,273]. 3. Other Mononitroarenes The products of the electrochemical reduction of mononitroarenes other than nitrobenzene are of the same kind as those reported earlier. Most studied are the nitronaphthalenes [284–290] for which the preparative electrochemistry can be summarized as follows. Under neutral conditions, reduction of 1- and 2-nitronaphthalene proceeds to the hydroxylamines, whereas reduction under acidic conditions at low potentials gives the corresponding amines. At more moderate reduction potentials, 1-nitronaphthalene gives 4-amino-1-naphthol, whereas 2-nitronaphthalene gives 2-amino-1naphthol [288]. Reduction of 1-nitronaphthalene at a Ti/TiO2 cathode under acidic conditions gives 1-aminonaphthalene [287]. Reduction at an amalgamated copper cathode in a sulfuric acid solution containing SnCl2 has been reported to give 1-amino-5-naphthol [285], whereas reduction in >50% H2SO4, MeOSO3H, MeSO3H containing an alcohol results in the formation of the corresponding 4-alkoxy-1-naphthylamines [286]. For instance, 2-methyl-1-nitronaphthalene could be converted to 4-methoxy-2-methyl-1-naphthylamine in 65% yield in this way. Reduction of the substituted nitronaphthalenes follows essentially the same rules as those that govern the products for the alkoxynitrobenzenes [289]. Heterocyclic compounds may result from 1-nitronaphthalenes substituted in 2- or 8-position and 2-nitronaphthalenes substituted in 1- or 3-position with CONH2, CN, and NHCOCH3 [290]. 4-Nitrobiphenyl is reduced at pH ≤ 9 in a 4F and a 2F step, presumably to the hydroxylamine and the amine [291]. 4. Di- and Trinitroarenes The mechanisms of the electrochemical reduction of di- and trinitrobenzenes (DNB and TNB) and other polynitroarenes are basically the same as those discussed earlier; still there are some

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characteristic features of the reaction patterns that originate from the presence of one or more additional nitro groups and that justify a separate section. The reactivity of the radical anions and dianions of, for instance, dinitrobenzenes is lower than that for the mononitrobenzenes, and this has allowed for studies of the kinetics of the protonation of the radical anions and dianions by using conventional electroanalytical methods. Also, the electronic coupling between the two nitro groups is dependent on their relative position. The 1,4-DNB dianion, for instance, is well described by a quinoid structure, whereas the 1,3-DNB dianion has biradical character as mentioned earlier and this affects the products of reduction. Finally, dinitrobenzenes carrying a substituent are most often unsymmetrical and this raises the question of which of the two nitro groups is first converted to a lower oxidation state by reduction. a. Mechanisms and Preparative Aspects of the Protonation in Aprotic Solvents The kinetics of protonation of the radical anions and dianions derived from the three simple dinitrobenzenes have been studied in DMF for a variety of proton donors including phenol and benzoic acid [41,43,292,293]. The rate constants for protonation of the radical anions by benzoic acid were k12 ═ 6.3·102 M−1 s−1, k13 ═ 3.6·102 M−1 s−1, and k14 ═ 670 M−1 s−1 [293], where the subscripts refer to the substitution pattern of the dinitrobenzenes, reflecting the general trend that the rate constants decrease in the order k12 > k13 >> k14 for both the radical anions and the dianions. Thus, the radical anions and dianions of 1,4-DNB are by far the least reactive. Reduction of 1,2-, 1,3- and 1,4-DNB in aprotic solvents in the presence of a suitable proton donor leads first to the corresponding nitrophenylhydroxylamines, possibly by a DISP-like mechanism [22,41,47,294]. Further reduction then leads to the phenylenediamines for 1,2-DNB and 1,4-DNB [22,47], illustrated in Equation 30.27 for 1,2-DNB, and to the benzenedihydroxylamine for 1,3-DNB [22,47] (Equation 30.28), illustrating that the two nitro groups in the latter case behave essentially independently [295]. NO2

NO2

+4e–,+4H+ –H2O

NH2

NHOH NO2

+4e–,+4H+ –H2O

NO2

(30.27)

–3H2O

NO2 NO2

NH2

+8e–,+8H+

NHOH

+4e–,+4H+

(30.28)

–H2O NHOH

NHOH

b. Mechanisms and Preparative Aspects of the Protonation in Water and Other Protic Solvents i. Unsubstituted Dinitrobenzenes and Dinitrobiphenyls The products resulting from electrochemical reduction of the three unsubstituted dinitrobenzenes [159,176,296–301] and of dinitrobiphenyls [291] in aqueous solution depend, as for the mononitrobenzenes, on pH of the solution. And as given earlier, the 1,3-isomer differs from the 1,2- and 1,4-isomers. Reduction of 1,2-DNB and 1,4-DNB at mercury or silver cathodes at low pH leads ultimately to the phenylenediamines [176,296,301] in overall 12F processes, but a 10F process leading first to the 1,2-quinonediimine that subsequently dimerizes to 5,10-dihydrophenazine-2,3-diamine and related compounds has been reported for 1,2-DNB as well [297]. The pathway to the phenylenediamines includes the 2- and 4-nitrophenylhydroxylamines [159,176,296] as in Equation 30.27, but may partly include also the azoxybenzene derivative [176]. In contrast, the final product resulting from reduction of 1,3-DNB at low pH is either 3-nitroaniline [300] or benzene-1,3-dihydroxylamine [296], in latter case via the 3-nitrophenylhydroxylamine as in Equation 30.28.

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Reduction of 1,2-DNB and 1,4-DNB at pH close to neutral gives the phenylenediamines [176,296] via the nitrophenylhydroxylamines [296] and the azoxy derivative [176]. The 1,3DNB gives 3-nitrophenylhydroxylamine [300] or benzene-1,3-dihydroxylamine [296] via the 3-nitrophenylhydroxylamine. Finally, reduction of 1,2-DNB and 1,4-DNB at high pH gives the phenylenediamines [176,296,299] or the nitrophenylhydroxylamines [298,302], whereas 1,3-DNB gives the benzene-1,3-dihydroxylamine [296] or 3,3′-dinitroazoxybenzene [298–300]. ii. Substituted Dinitrobenzenes and Dinitrobiphenyls The electrochemical reduction of substituted dinitrobenzenes has been investigated carefully [86,303–310], and the effect of substituents on the reduction potentials has been reviewed [311]. Reduction of monosubstituted 1,2- and 1,4-dinitrobenzenes at a mercury cathode under acidic conditions leads first to the nitrophenylhydroxylamines in a 4F process and then at a lower potential to the 1,2-phenylenediamines in a 12F process [303–305]. At Ti/TiO2 cathodes the direct reduction to the diamines is observed [204,312,313] similar to what was described to mononitrobenzenes earlier. For 1,2-dinitrobenzenes (X ═ 3-Me, 3-OEt, 3-Br, 3-COOH, 4-Me, 4-OMe, 4-Cl), the 4F process involves the nitro group situated meta to X. Only in one case (X ═ 4-COOH) the nitro group in the para-position relative to X is being reduced [303]. For the related 1,4-dinitrobenzenes, the 4F reduction of the nitro group ortho to X is involved for electron-withdrawing substituents (X ═ 2-Br, 2-COOH, 2-CONH2), whereas the nitro group meta to X is being reduced for electron-donating substituents (X ═ 2-Me, 2-OH, 2-OMe) [304]. The substituted 1,3-dinitrobenzenes behave differently. Except for X ═ OR the reduction includes three successive 4F processes, first to the nitrophenylhydroxylamines, then to the benzenedihydroxylamines and finally to phenylenediamines [308]. Also in this case, the nature and position of X affects the selectivity of the reduction. For the halides (X ═ 4-Cl, 4-Br, 4-I), the nitro group ortho to X is first reduced, but for alkyl and carboxy substituents (X ═ 4-Me, 4-Et, 4-COOH), the nitro group meta to X is first reduced. When X ═ 2-OR or 4-OR, two 6F processes are observed; the first that includes hydrolysis of the OR substituent leads to 2-amino-4-nitrophenol and the second to the 2,4-diaminophenol. The reduction of 2,6-dinitrotoluene follows the pattern outlined earlier for nitrobenzenes. Under strongly acidic conditions (sulfuric acid), reduction leads to 2-methyl-3-nitroaniline, under weakly acidic conditions (acetic acid/acetate buffer) the product is 2-methyl-3-nitrophenylhydroxylamine and under slightly basic conditions (acetate in ethanol) 2,2′-dimethyl-3,3′-dinitroazoxybenzene is produced [306]. Reduction of 2,4-dinitrotoluene [306], 2,4-dinitroanisole [309] and 2,4-dinitrochlorobenzene [309] proceeds in essentially the same way. Systematic studies of the conversion of 2,4-dinitrophenol [307] and 3,5-dinitro-2-methylphenol [310] to the corresponding diamines have resulted in conditions for high yield processes (~80%). An interesting variant is reported for the reduction of 2,3′-dinitrobenzidine that in buffered aqueous methanol at all pH values is reduced in a 10F (4F + 4F + 2F) process to 2-amino-3′hydroxylaminobenzidine via a quinoid intermediate (Equation 30.29) [314], illustrating the effect of having at same time a nitro group ortho and meta to the amino group in the starting material. H2N

NH2 NO2

2.(+4e–,+4H+) 2.(–H2O)

H2N

NH2 NHOH

NO2

NHOH

(30.29) +2e–,+2H+ –H2O

NH

H2N NH

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NHOH

H2N

NH2 NH2

NHOH

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iii. Dinitronaphthalenes A series of dinitronaphthalenes has been studied by polarography, coulometry, and preparative electrolysis [284,288,315]. Two or three reduction processes are observed under acidic or close-to-neutral conditions, the number of polarographic waves and the number of electrons involved in each step being dependent on the substitution pattern. The behavior of the 1,2- and 1,4-isomers is similar to that of 1,2-DNB and 1,4-DNB, that is, the reduction proceeds as a 4F process to the nitronaphtylhydroxylamine and as a 12F process to the naphthalenediamines. The remaining isomers (1,3; 1,5; 1,6; 1,7; 2,3; 2,6; and 2,7), except for 1,8-isomer, give rise to three waves corresponding to 4, 8, and 12F, respectively, with the 8F processes most likely corresponding to the formation of the unstable naphthalenedihydroxylamines. Preparative electrolysis in aqueous ammonium acetate/ethanol results in the formation of the nitronaphthylhydroxylamines in a 4F process and the general trend observed is that the 1-nitro group is the one being reduced in the unsymmetrically substituted compounds. Reduction under acidic conditions leads in all cases to the naphthalenediamines in 12F processes. The 1,8-isomer is a special case. This compound and related 4,5-disubstituted derivatives are reduced in 8, 10, or 12F steps with the number of electrons decreasing with increasing pH [316] corresponding to the formation of benzo[cd]indazole (1,8-naphthopyrazol, the 8F intramolecular azo compound), 1,2-dihydrobenzo[cd]indazole (1,8-naphthodihydropyrazol, the 10F intramolecular hydrazo compound), and naphthalene-1,8-diamine (the 12F product) (Scheme 30.9). Electrosynthesis of naphthalene-1,8-diamine at a large scale may be carried out in a close-toquantitative yield by reduction of a suspension of 1,8-dinitronaphthalene at 70°C at a copper cathode in aqueous sulfuric acid containing a titanium salt [317]. c. Other Reactions of Di- and Trinitroarenes Elimination of a nitro group as a nitrite ion has been reported for the reduction of 1,4-DNB [318] and of 1,2,4,5-tetrafluoro-3,5-dinitrobenzene in DMF containing a proton donor [319]. These reactions are similar to what is observed for aliphatic nitro compounds (see Sections II.B.3.b.i and II.B.4.a), but we are not aware of other examples of this type of cleavage for an aromatic nitro compound. Another unusual reaction, in DMF in the presence of N-methylformamide, is the reductively induced substitution of a hydrogen atom in 1,3,5-TNB with an N-methylformamido group to give N-methyl-N-(2,4,6-trinitrophenyl)formamide in 78% yield [320]. The first step appears to be the reaction between the 1,3,5-TNB radical anion and N-methylformamide to form a Meisenheimertype adduct. N

N

+8e–,+8H+ NO2

NO2

–4H2O

HN

NH

+10e–,+10H+ –4H2O +12e–,+12H+

NH2

NH2

–4H2O

SCHEME 30.9

The 8, 10 and 12F products resulting from the reduction of 1,8-dinitronaphthalene.

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5. Hetereoaromatic and Related Mono- and Dinitro Compounds This section includes compounds in which the nitro group is attached directly to a heteroaromatic ring. Basically, heteroaromatic nitro compounds behave in much the same way as nitroarenes during electrochemical reduction. However, occasionally the properties of the heteroaromatic ring add to the complexity typically observed for the reduction of simple nitroarenes. For example, if the nitro group is attached to the heteroatom, X, reduction often results in cleavage of the X–NO2 bond. Other examples include heteroaromatic systems that are inherently acidic owing to the presence of, for example, an N–H group in the ring; this may lead to self-protonation reactions similar to those described earlier for nitrophenols and nitrobenzoic acids. Finally, the heterocyclic chemistry is rich on positively charged rings such as pyrazolium and imidazolium that have only few equivalents in arene chemistry. a. Compounds with the Nitro Group Attached to Carbon i. Reduction in Aprotic Solvents 1-Methyl-4-nitroimidazoles and 1-methyl-5-nitroimidazoles are reduced in DMF to the corresponding radical anions in contrast to the unsubstituted 4- and 5-nitroimidazoles that owing to the acidic 1-NH proton are reduced to the corresponding hydroxylamines in a self-protonation process [63,64,67]. Similarly, self-protonations have been observed in DMSO for 4-substituted 7-nitro-3,4-dihydroquinoxalin derivatives and 5-nitro-2,3-dihydro-1Hindazole derivatives [62]. ii. Reduction in Water and Other Protic Solvents The electrochemical reduction of this class of compounds in aqueous solution follows closely that of the simple nitroarenes although with the additional complexity that protonation of the heterocyclic ring may take place at low pH. Thus, 4-nitropyridine [154] and 4-nitropyridine-N-oxide [154,155] as well as nitrofurans [68], nitropyrazoles, nitroimidazoles, and nitrotriazoles [63,64,321–328] are all reduced to the corresponding hydroxylamines that at low pH may be further reduced to the amines [321,322] or oxidized to the nitroso compound [328]. Under basic conditions, it is observed that the hydroxylamines resulting from reduction of 2-alkyl-4-nitroimidazoles and 1,2-dialkyl-5-nitroimidazoles eliminate water easily to give the quinone imine that is further reduced to the amine in an overall 6F process [322]. Similarly, the 1-alkyl4-amino-5-nitroimidazole and 1-alkyl-4-nitro-5-aminoimidazole are reduced in 6F processes [322]. In contrast, the reduction of 5-nitro-1,2,4-triazol-3-one in aqueous sulfuric acid leads to the azoxy derivative as the major product that owing to its low solubility can be isolated by filtration [327,329]. The reduction of nitropyridines and the hydroxy- and methoxy-substituted derivatives in protic solvents follows essentially the same rules as discussed earlier for the analogous nitrobenzenes [158,330]. For 4-nitropyridine, the +2e−, +2H+ reduction to the dihydroxyamine was found to be reversible at −7°C in acidic ethanol [158]. b. Compounds with the Nitro Group Attached to a Heteroatom The reduction of 1-nitropyrazole in acidic media results in cleavage of the N–N bond and the formation of nitrous acid [331] and thus resembles the reduction of nitramines (see Section V.A). At pH > 4, competitive reduction to the nitrosamine takes place. c. Positively Charged Compounds The nitro group attached to a 1,2-dialkylpyrazolium or 1,2,3-trialkylimidazolium ring is reduced preferentially to the positively charged heterocyclic system at pH up to approximately 9. Reduction leads to the corresponding hydroxylamines with the sequence of the electron and proton transfer steps, leading to the intermediate nitroso derivative being H+, e−, e−, H+ [331].

B. ALIPHATIC NITRO COMPOUNDS The electrochemical reduction of compounds in which the nitro group is attached to an sp3-hybridized carbon (nitroalkanes and related) or a nonaromatic sp2-hybridized carbon (nitroalkenes and related)

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1171

is discussed in this section. We are not aware of reports addressing the electrochemical reduction of compounds in which the nitro group is attached to an sp-hybridized carbon (nitroalkynes and nitroallenes) in spite of the fact that such compounds are known. Aliphatic nitro compounds in which an aromatic substituent is attached to a nonaromatic carbon are included in this section as well. 1. Formation and Properties of the Radical Anions The half-life of the radical anions derived from nitroalkanes is as a rule shorter than that of radical anions derived from aromatic nitro compounds, typically in the range, 0.1 < t ½/s−1 < 1. The radical anions derived from primary and secondary nitroalkanes have been observed in aprotic solvents such as MeCN by fast scan CV [101,332] or in high-speed channel cells [333]. Tertiary nitroalkanes that lack the acidic α-hydrogens are less reactive and the one-electron reductions have been studied in detail by slow sweep CV in MeCN [30,92,101,332,334–343], DMF [102,337,344–346], DMSO [337,347–349], pyridine [337], 1,2-dimethoxyethane (glyme) [92,339,349,350], and room temperature ionic liquids [29,30]. Primary and secondary nitroalkanes exist as the nonreducible anions under basic aqueous conditions, but deprotonation is usually slow and it has occasionally been possible to observe the oneelectron reduction in basic aqueous solutions [351]. The radical anions of tertiary nitroalkanes have been observed in aqueous solution as well [352]. Radical anions have been observed also for di- and trinitroalkanes in DMF [344], glyme [350], MeOH [353] and aqueous base [352,354] and for nitroalkenes in MeCN [355,356] and DMF [357], but geminal dinitroalkanes are inherently unstable owing to the facile loss of a nitrite ion (see Section II.B.4). Reduction potentials (half-wave potentials) for a series of aliphatic primary and secondary mono- and dinitro compounds [353] and tertiary nitro compounds [334,350] have been recorded in various nonaqueous solvents and solvent mixtures and the dependence of ion-pair formation on the reduction process has been discussed [92,332]. The kinetics of the heterogeneous electron transfer process for reduction of aliphatic nitro compounds have been studied extensively [30,101,102,332,333,336–338,341–343,346–349]. In particular, 2-methyl-2-nitropropane has featured as a prototype compound in such studies [30,101,102,332,336,337,341–343,346–349]; this compound is one of the few examples of organic substrates for which the electron transfer process appears as quasireversible even at slow sweep CV and at the same time the radical anion has an appreciable lifetime. The values of ks have been determined by CV and by impedance measurements and are typically in the range 10 –4 < ks/(cm s−1) < 10 −1 depending particularly on the size of the electrolyte cation (R4N+) [30,101,102,332,333,336–338, 341–343,346–349]. In general, the value of ks decreases with increasing size of R4N+, which is likely to reflect different planes of closest approach [336]. The effect of solvent appears to be smaller than predicted by the Marcus theory [337]; this may be caused by dielectric saturation effects and/or a solvent-dependent frequency factor, and it is most likely that also ion-pair effects that vary with the solvent are important in determining ks [101,332,336–338,348,349]. The value of the transfer coefficient, α, is typically around 0.45 and varies with the potential as predicted by the Marcus theory; typically it is found that dα/dE is in the range 0.25–0.5 V−1 [332,336,338,348,349]. As for nitrobenzene, it is found that the diffusion coefficients in ionic liquids, ∼10 −7 cm2 s−1, are orders of magnitude smaller than in conventional aprotic solvents [29,30]. 2-Methyl-2-nitropropane has been used to evaluate the predictions made by the Marcus–Hush and the Butler–Volmer kinetic formalisms (see Chapter 1 for details), and it was found that the electrochemical data could be fitted satisfactorily to both models provided that the asymmetric Marcus–Hush model was used [333,342,343,347]. The quasireversible electron transfer for 2-methyl-2-nitropropane may be catalyzed by a redox mediator such as terephthalonitrile [332,338]. The rate constant for electron transfer from the radical anion of terephthalonitrile to 2-methyl-2-nitropropane in MeCN was found to be 5·105 M−1 s−1 and the

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self-exchange rate constant for 2-methyl-2-nitropropane was 5·103 M−1 s−1. The temperature dependence gave access to the activation energy that was found to be 25–30 kJ mol−1. Intramolecular redox catalysis has been studied also [358,359]. Finally, it should be mentioned that myoglobin has been applied as a redox catalyst for nitromethane reduction [360]. The lifetime of the radical anions of many nitroalkanes and nitroalkenes is sufficiently long to allow for recording of the ESR spectra [174,335,339,350,355,357,361]. In contrast to what is observed for nitrobenzenes, the unpaired electron has been observed to be almost exclusively localized at the nitro group. 2. Nitroalkenes and Related a. Simple α,β-Unsaturated Nitro Compounds i. Reduction in Aprotic Solvents Reduction of α,β-unsaturated nitro compounds in aprotic solvents leads to the radical anions that subsequently dimerize. The resulting dimer dianions may be converted by protonation to the corresponding 1,4-dinitroalkanes in good yields [362–365]. The dimerization may also be carried out electrocatalytically; in this case, 1,3-dinitrodimers are formed [363]. A more thorough reduction occurs when the reaction is carried out at a carbon cathode in DMF in the presence of TiCl4 and with Et4NTos as supporting electrolyte. In that case, nitriles are formed in good yields (Equation 30.30) [366]. This is a convenient method for the conversion of aldehydes to one-carbon elongated nitriles through the preparation of the nitroalkenes by condensation of the aldehyde with nitromethane. Ar

CH

CH

NO2

5–6 F 64–95%

Ar

CH2CN

(30.30)

ii. Reduction in Water and Other Protic Solvents The reduction of α,β-unsaturated nitro compounds under acidic conditions in the presence of an organic co-solvent leads first to the oxime [362,367–371] in 40–90% yield, presumably via the enehydroxylamine (Equation 30.31). One of the by-products, resulting from hydrolysis, is the corresponding ketone that in a separate step may be converted to the oxime by reaction with hydroxylamine resulting in an overall yield of the oxime better than 90% [362]. R1

R3 C

+2e–,+2H+

R1

R3 C

C

+2e–,+2H+

C

–H2O R2

NO2

R1

R3

NO

R2

(30.31) C R2

R1

C

R3 CH

NHOH

R2

C NOH

The oxime may be further reduced to the amine [362,371]. Similar to the preparation of chainelongated nitriles reported earlier, these reactions are employed for the synthesis of oximes and carbonyl compounds with a longer carbon chain. Reduction to the amine state has been put to use in the preparation of amino acids, for instance, tryptophan ethyl ester (50–60%) by reduction of ethyl 3-(3-indolyl)-2-nitroacrylate [372]. 2,3-Dinitro-2-butene gives rise to a 2F polarographic wave at all pH values followed in acidic solution by another pH-dependent reduction wave [373,374]. Preparative reduction in acidic solution results in complex reaction mixtures containing 2-nitro-2-butene and diacetylmonoxime and also hydrolysis products of the latter; the former is the major product at higher pH. Reduction of

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the conjugated 1,4-dinitro-1,3-butadiene at pH 1.4 results in partial saturation of the butadiene to give a 74% yield of 1,4-dinitro-2-butene [375]. b. Halogen-Substituted α,β-Unsaturated Nitro Compounds When speaking here about α,β-unsaturated nitro compounds, the Greek letters refer to the position of the double bond relative to the nitro group. However, in the conventional nomenclature for compounds of this type, the assignment of the carbons is reversed. Thus, β-bromo-β-nitrostyrene in the present context is an α,β-unsaturated nitro compound. Reduction of β-bromo-β-nitrostyrene in acidic solution gives benzyl cyanide (80%) (Equation 30.32) and benzaldehyde (10%) [373,376]; the related 1-bromo-3-methyl-1-nitro-1-butene, in which the phenyl group is replaced by an iso-propyl group, gives a complicated reaction mixture of various nitriles suggested to arise via an allene-type oxime that after elimination of water gives a reactive carbocation intermediate (Equation 30.33) [377]. Br Ph

CH

+6e–,+5H+

C

CH2CN + Br– + 2H2O

Ph

(30.32)

NO2 H3C

Br CH CH

H3C

+4e–,+3H+

C

CH CH NO2

H3C

H3C

C

NOH

+ Br– + H2O

+H+ –H2O

H3C

(30.33)

+

CH CH

Various nitriles

CN

H3C

α-Chloro-β-nitrostyrene gives a mixture of products including Ph-CHOHCHO that was believed to arise from hydrolysis of the oxime initially formed by reduction (Equation 30.34) [375] similar to the reaction for the unsubstituted compound (Equation 30.31). Ph CH NO2

C

+4e–,+4H+

Ph CH CH NOH + H2O

(30.34)

Cl

Cl

c. Nitrohydrazones Reduction of nitroacetaldehyde phenylhydrazone in 0.05 N sulfuric acid in 25% dioxane consumed 3.7F and the corresponding hydrazooxime, resulting from tautomerization of the initially formed azooxime, was detected as the product (Equation 30.35) [378]. CH3 Ph

N

NH

NO2

Taut.

CH3

+4e–,+4H+ –H2O

Ph

NH

N NHOH

(30.35) CH3

Ph

NH

NH NOH

Related to this is the 6F reduction of o-nitrophenylazo phenylnitromethane to o-nitrophenylbenzamidrazone that in turn may be converted to 3-phenylbenzo-1,2,4-triazine [379] reminiscent of the reduction benzamidoxime to benzamidine [380].

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3. Mononitroalkanes and Related Compounds that have a functional group attached to the molecule without being in electronic conjugation with the nitro group as a nitroalkane such as Ph–CO–CH2–NO2 and EtO–CO–CH2–NO2 are included in this section. Nitroalkanes that have potential leaving groups attached are treated in Section II.B.7. a. Primary and Secondary Nitroalkanes i. Reduction in Aprotic Solvents Usually nitroalkanes such as nitromethane cannot be reduced to the corresponding nitrosoalkanes by reduction in aprotic solvents. However, it has been observed that the trans-dimer of nitrosomethane may be obtained by reduction of nitromethane in 1-butyl3-methylimidazolium tetrafluoroborate on a copper disk cathode [381]. The dimers of the nitroso compounds are more difficult to reduce than the monomers, and this may be the reason why the nitroso compounds survive in this case. Reduction of weak acids to hydrogen gas and the corresponding anion is a common method for the preparation of a so-called electrogenerated base (see Chapter 43). However, in the case of nitro compounds such as ethyl nitroacetate, reduction is accompanied by considerable C–N cleavage. Instead, the nearly quantitative conversion to the anion was found when reduction was carried out in air-saturated solutions; the electrogenerated superoxide ion then serves as a base that effectively deprotonates the substrate (Equation 30.36) [382]. Only the first step is shown. NO2

NO2 H2C

+ O2 COOEt

–

HC

+ HO2

(30.36)

COOEt

The anion shown and also those derived from nitromethane, ethyl 2-nitropropionate, diethyl 2-nitroglutarate, 2-nitropropane, nitrocyclopentane, ethyl 2-nitro-4-pentenoate, and dinitromethane, have all been prepared in a similar fashion and have many applications in organic synthesis [383]. Similarly, the electrochemical reduction of a mixture of, for instance, an aldehyde in nitromethane as the solvent and with a nickel cathode leads to the formation of the nitromethane anion that in turn adds to the aldehyde in a Michael-type reaction [384–386]. The addition products are formed in isolated yields in the range 60–90%. ii. Reduction in Water and Other Protic Solvents Primary and secondary nitro compounds are weakly acidic (pKa ═ 10.2 for nitromethane) and exist as the nonreducible anions under basic conditions. The details of the first electron and proton transfer steps have been studied by polarography [387], impedance measurements [388–392], and application of channel electrodes [393,394], and the data have been analyzed in terms of an ECCe-type mechanism. Results from earlier work were in agreement with the formation of a radical anion as the first step [395]. Under acidic or neutral conditions, electrochemical reduction at room temperature leads to the corresponding N-alkylhydroxylamines in 60–90% yield [396–402], for instance, isolated as the hydrochloride [398], together with minor amounts of the N-alkylamine and a carbonyl compound. It has been shown [399] that the N-alkylhydroxylamines are not further reduced under the conditions and the origin of the N-alkylamine and the carbonyl compound was suggested to arise instead via tautomerization of the intermediate nitrosoalkane to the oxime and hydrolysis of the latter (Scheme 30.10). Here it is assumed that the electroactive species is the protonated [161] or at least strongly hydrogen-bonded nitroalkane. The reduction at mercury electrodes involves most likely adsorbed species [161,403]. The product from reduction of nitromethane, N-methylhydroxylamine, is an important compound in organic synthesis and detailed descriptions for carrying out the electrosynthesis of this compound at a large (industrial) scale have been reported [404,405]. In an interesting variant,

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1175

Reduction of Nitro Compounds and Related Substrates

HCR΄R

NO2H+

+2e–,+2H+ HCR΄R

NOH+

+2e–,+2H+ HCR΄R

–H2O

CR΄R

NOH2+

+4e–,+4H + –H2 O

NH2OH+

RR΄CHNH3+

H2O CR΄R

SCHEME 30.10

O + H3NOH+

Reduction of nitroalkanes to the oxime and hydroxylamine in aqueous solution.

the focus was on the formation of the oxime, and it was shown [406] that reduction of α-nitrobenzylic compounds in aqueous acetic acid buffers mixed with ethanol resulted in a mixture of the oxime and the hydroxylamine; the latter could then be oxidized, chemically or electrochemically, to the oxime via the nitrosoderivative. The overall yields of the oximes were in the range 80–90%. As seen from Scheme 30.10, the tautomerization of the nitrosoalkane competes with the further reduction to the N-alkylhydroxylamine and accordingly it would be expected that lowering the temperature would slow down tautomerization and thus favor the further reduction. In contrast, increasing temperature would be expected to favor the formation of the N-alkylamine. This is indeed what is observed and it has been reported that, for instance, the electrochemical reduction of α-hydroxynitroalkanes at 65–86°C leads preferentially to the aminoalcohols, whereas hydroxylaminoalcohols are the products when the reduction is carried out at 0–35°C [407]. Also, the rate of tautomerization of the nitrosoalkane to the oxime depends on pH and on the nature of the substituents, R and R′ and, as the rule, the rate increases with increasing electronegativity of R and R′ [408–413] and, accordingly, a higher proportion of the amine is observed when one or both of R and R′ are strongly electron withdrawing. In the extreme, the amine is the only product as it has been observed during the reduction of 2-nitromalonic ester (R ═ R′ ═ COOEt) to 2-aminomalonic ester at low pH [412,414]. The hydroxyiminomalonate that is an intermediate in this reaction may be reduced electrocatalytically to the 2-aminomalonate at a Ti/nanoporous TiO2 electrode [415]. Similar to what is observed for aromatic nitro compounds, the formation of heterocyclic products is observed when the nitro group is positioned strategically relative to another functional group. The reduction of γ-nitroesters to N-hydroxypyrrolidinones [401] serves to illustrate this aspect (see Chapter 34 for a general discussion of the formation of heterocyclic compounds). b. Tertiary Mononitroalkanes i. Reduction in Aprotic Solvents The radical anion of 2-methyl-2-nitropropane has a life time, t½, of approximately 0.66 s in glyme (1,2-dimethoxyethane) [350] and decomposes into nitrite ion and a tert-butyl radical [334,335,339,346,350,416]. When generated electrochemically, continued electrolysis was found to give di-tert-butylnitroxide, ((CH 3 )3 C)2 N−O• [335,339,350]. However, other products may arise from the follow-up reactions of electrogenerated tert-alkyl radicals depending on whether the radical is further reduced under the reaction conditions and on the lifetime of the radical anion. Some of these pathways are summarized in Scheme 30.11. Dimerization of the free radicals, R3C•, was observed for tert-nitrocumene [350] and 2-(4-nitrophenyl)-2-nitropropane (Scheme 30.12) [345]. In the latter case, the radical anion was found to cleave heterolytically as indicated [340]. Complete reduction to the alkane, and other products, was observed for 2-nitro-2,4,4-trimethylpentane [350]. When 2-methyl-2-nitropropane is reduced in MeCN in the presence of a proton donor such as PhOH, protonation competes effectively with elimination of nitrite ion and the hydroxylamine is formed [334].

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1176

Organic Electrochemistry +e– R3C

NO2

–e–

R3C R3C

R3C

R3C NO2 –

R3C

slow

NO2 –

R 3C

+ NO2–

R3C

NO + R3CO–

CR3 (R3C)2NO2–

+e– R3C

R3C–

(R3C)2N

O

R3C

SCHEME 30.11

OCR3

Pathways for the reduction of tert-nitroalkanes. –

CH3 O2N

CH3 NO2

–NO2

–

O2N CH3

CH3

CH3 CH3

CH3 O2N

2O2N CH3

SCHEME 30.12

(R3C)2N

NO2 CH3 CH3

Mechanism for the reduction of 2-(4-nitrophenyl)-2-nitropropane.

ii. Reduction in Water and Other Protic Solvents The tertiary nitroalkanes differ from the primary and secondary compounds in that the intermediate nitroso derivatives cannot rearrange into the oximes. Accordingly, the yield of alkylhydroxylamine is usually high (80–90%) [399]. The nitroso intermediate may in this case be detected directly in the reduction mixture by its bluish green color and its polarographic wave. The hydroxylamines are reduced to the corresponding amines at a lower potential [417] and resemble in that respects the aromatic nitro compounds. Similar to what was observed in aprotic media, it was found that tert-nitrocumene could be reduced at mercury in basic aqueous methanol to the R3C–CR3 dimer; however, if instead a porous Raney nickel cathode was used, catalytic hydrogenation resulting in the formation of the amine took place almost exclusively [418]. 4. Dinitroalkanes a. Reduction in Aprotic Solvents The radical anions resulting from the reduction of nonacidic geminal nitroalkanes such as 2,2-dinitropropane [120,419] and 1,1-dinitrocyclohexane [344,354,420,421] in DMF undergo cleavage of a C–N bond forming a nitroalkyl radical and nitrite ion. The following steps for the reduction of 1,1-dinitrocyclohexane leading to the formation of the vicinal 1,1′-dinitrobicyclohexyl include a chain reaction between the nitronate anion and the starting material (Scheme 30.13). A minor product is nitrocyclohexane resulting from further reaction of the nitrocyclohexyl radical. Related compounds in which one of the nitro groups is replaced by another electron-withdrawing group react similarly [344]. Vicinal dinitro compounds such as 1,1′-dinitrobicyclohexyl (the product given earlier) have been investigated carefully [420]. Reduction in DMF, here illustrated by 2,3-dinitro-2,3-dimethylbutane, proceeds according to Scheme 30.14. The rate constant for elimination of nitrite ion from the radical anion was close to 103 s−1.

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1177

Reduction of Nitro Compounds and Related Substrates NO2

NO2 –

+e– –e–

NO2

NO2

NO2

–NO2–

+e– NO2

NO2 –

+

NO2–

NO2 NO2 – NO2

– NO2

NO2

NO2

+

+ NO2

SCHEME 30.13 H3C

O2N

Mechanism for the reduction of 1,1-dinitrocyclohexane. CH3 NO2

O2N H3C

NO2–

H3C

+e–

O2N

–e–

CH3

CH3 –

H3C

NO2

H3C

–NO2–

CH3

O2N

CH3

H3C

CH3

+e– H3C

CH3

O2N H3C

H3C O2N H3C

SCHEME 30.14

CH3

CH

–

H3C + O2N H3C

CH3 NO2 CH3

H3C O2N H3C

CH3 NO2 CH3

–

+e– NO2 CH3

H3C

CH3 +

H3C

NO2–

CH3

H3C

CH3

H3C

CH3

–2NO2–

Mechanism for the reduction of 2,3-dinitro-2,3-dimethylbutane.

b. Reduction in Water and Other Protic Solvents The reduction of geminal dinitroalkanes carrying an acidic hydrogen in the α-position depends on both the structure and not least on pH of the electrolyte solution and related to this, the position of the equilibrium between the nitro and aci-forms. An additional complication is that the aci-form may suffer hydrolysis to the corresponding carbonyl compound under acidic conditions (the Nef reaction). The reduction of 1,1-dinitroethane in acidic solution required 5–6F and nitrous acid was detected in the resulting solution [422]. The following sequence of steps (Equation 30.37) leading to the hydroxamic acid oxime via the intermediate formation of the nitrolic acid has been proposed [423–425]. The reduction of dinitromethane appears to follow a similar route, but is further complicated by nitrosation of the starting material with the nitrous acid liberated during reduction [425].

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1178

Organic Electrochemistry

H3C H

NO2

O–

H3C

+2e–

N+

NO2

+

N+

O–

H

OH

H3C

+2H+

NO2–

+

HNO2

O–

H

(30.37) NOH –H2O

H3C

NOH

+4e–,+4H+

C

H3C

–H2O

NO2

C NHOH

The anions resulting from deprotonation of acidic 1,1-dinitroalkanes are reduced in a reversible or quasireversible one-electron process to ion-paired dianion radicals that suffer protonation and further reduction to the corresponding nitrolic acids and hydroxamic acid oximes [426–428] or cleavage of the C–N bond [429] depending on the conditions. The anions have a high affinity for mercury and give rise to an anodic polarographic wave resulting from the formation of organomercurials [430]. The ESR spectra of the dianion radicals have been recorded in a number of cases [361,431]. In slightly acidic solution, 2,2-dinitropropane is first reduced to the nitronate anion and nitrite in a 2F process; further reduction of the resulting pseudonitrole leads to acetone oxime (Equation 30.38) [422,432,433]. H3C H3C

NO2

O–

H3C

+2e–

N+ H3C

NO2

+

+2H+

NO2–

O–

H3C

NO

H3C

NO2

(30.38) H3C

+2e–,+2H+

NOH +

HNO2

H3C

The formation of the pseudonitrole requires the presence of the nitrous acid also formed in the first step. If the nitrous acid is removed by adding a scavenger or by reduction (to hydroxylamine and ammonia), which may happen at lower potentials, the nitronate anion is instead protonated to 2-nitropropane that in turn is reduced to the 2-hydroxylamine [419]. In basic solution, reduction stops after the formation of the 2-nitropropane anion and nitrite (Equation 30.39) [422,432,433]. Upon addition of acid to the resulting solution, the corresponding pseudonitrole is formed as earlier. H3C H3C

NO2

+2e–

O–

H3C

N+ H3C

NO2

(30.39)

NO2–

+

O–

Nitrolic acids are reduced in acidic solution to N-hydroxyacetamidoximes [434,435] and further to N′-hydroxyacetimidamide and acetamidine [435]; the N-hydroxyacetamidoximes may be oxidized anodically to the corresponding nitrosolic acids. In alkaline solution, nitrolic acids have been reported to undergo reductive dimerization to azo derivatives (Equation 30.40) [436]. 2H3C

NO2

+8e–,+8H+

NOH

–4H2O

C

H3C

C

N

NOH

N

C

CH3

NOH

(30.40)

5. Trinitroalkanes The reduction of trinitroalkanes [425,437–439] proceeds similarly to reduction of dinitroalkanes; however, the presence of the extra nitro group causes the intermediate nitronitrolic acid (dinitroformoxime)

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1179

Reduction of Nitro Compounds and Related Substrates

[440] derived from trinitromethane to be reduced to 1,3-dihydroxyguanidine [425,437,439]. In base, the 1,3-dihydroxyguanidine is slowly converted to nitrosoformimidamide that in turn is reduced to 1-hydroxyguanidine [437]. The first step in the reduction of 1,1,1-trinitroethane includes as expected from the pattern outlined earlier reductive cleavage to the dinitromethane anion and nitrite [361,441]. The anion may then undergo nitrosation, protonation, and/or the Nef reaction [441]. The reduction of the anions leads to elimination of nitrite [429], but detailed polarographic studies are hampered by adsorption [442]. 6. Tetranitroalkanes The tetranitromethane radical anion undergoes cleavage to the trinitromethane anion and nitrite ion as expected [395,437,438,443]. The yield of trinitromethane is high (80%) when reduction is carried out at a platinum electrode in acidic solution [438], but drops to around 45% in neutral solution. A parallel, nonelectrochemical reaction between tetranitromethane and mercury has been observed; the details of this reaction are not clear [437]. 7. α-Halonitroalkanes Nitroalkanes with a potential leaving group in the α-position are observed to undergo elimination upon electrochemical reduction. Typically the leaving group is a halide ion (Cl−, Br−, or I−; see also Chapters 24 and 25) but may also be other good leaving groups such as p-toluenesulfinate [344] or NO2− as in the reduction of 1-cyano-l-nitrocyclohexane [344] and in the geminal dinitroalkanes discussed earlier. This raises the question of which anion is eliminated during the reduction of an aliphatic compound having a nitro group and a potential leaving group attached to the same carbon. Using the halides as an example, the first step of the reductions may proceed according to one of the following two 2F routes (Equations 30.41 and 30.42). The resulting carbanion may undergo a variety of reactions. The abbreviations Y and Z symbolize a hydrogen atom, an alkyl group, an additional halogen atom (Hal), or an additional nitro group. Y

Y +2e–

Hal

O

NO2

+

Z

+2e– NO2 Z

Z

(30.41)

O–

Y

Y Hal

+ Hal–

N

N O–

Z

O–

Y

+ –

–

Hal

+ NO2–

(30.42)

Z

Reduction of α-monohalo-α-mononitroalkanes (Hal ═ Cl, Br, I; Y ≠ NO2; Z ≠ NO2) proceeds according to Equation 30.41 with the loss of Hal− and the formation of the nitronate ion that subsequently may undergo a variety of reactions, including protonation and further reduction, depending on the structure of the substrate, the experimental conditions, and the addition of suitable reactants to the electrolysis solution [423,444–446]. The reaction is governed by the stability of the resonance stabilized nitronate anion that may gain additional stability by, for instance, the presence of a neighboring carbonyl group as in 2-bromo-2-nitroacetophenone [400] and in α-halo α-nitro derivatives of camphor [447]. Similarly, the reduction of α,α-dihalo-α-mononitroalkanes [423] and α,α,α-trihalonitromethanes [423,448] follows Equation 30.41; in those case, the halogen atoms may be lost one-by-one in altogether two or three reduction, elimination, protonation sequences. For example, reduction of 2-halo-2-nitropropanes (Y ═ Z ═ Me) in aprotic media results in the formation of 2,3-dimethyl-2,3-dinitrobutane in good yields [444]. The reaction was suggested to most likely include a 2F reduction to the 2-nitropropane anion (Equation 30.41) that subsequently reacts with the starting material in nucleophilic substitution reaction. Similarly, reduction of 3,7-dibromo-3,7-dinitrobicyclo[3.3.1]nonane in

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1180

Organic Electrochemistry

aqueous acetone gave 3,7-dinitronoradamantane in 75–80% yield [446]. Related to these reactions is the reduction of 2-bromo-2-nitropropane at mercury in dichloromethane in the presence of methyl propiolate that gives a mixture of 2,3-dimethyl-2,3-dinitrobutane, 2-nitropropane, methyl (E)-4methyl-4-nitro-2-pentenoate and an organomercurial. A more unusual reaction was observed in the presence of dimethyl acetylenedicarboxylate from which the two products shown in Equation 30.43 could be identified [445]: H3C

NO2

H 3 COOC

+

O–

H3C

H3COOC

H3COOC CHCOOCH3

+

(30.43)

COOCH3

H3C

H3C O2N

COOC H 3

N

–Br–

Br

H3C

O–

H3C

+2e–

CH3

H3C

COOCH3

N

A variety of different products also resulted from the reduction of 5-halo-5-nitro-1,3-dioxanes and 2-halo-2-nitro-1,3-propanediols and also in this case the first reaction proceeds according to Equation 30.41 [449]. The first reduction step for α-monohalo-α,α-dinitroalkanes (Hal ═ Cl, Br, I; Y ═ NO2; Z ≠ NO2) [441] and α-monohalotrinitromethanes (Hal ═ Cl, Br, I; Y ═ Z ═ NO2) [450] depends both on the halogen and the number of nitro groups. For example, during the reduction of α-bromo-α,α-dinitroethane, NO2− is eliminated (Equation 30.42), whereas for bromotrinitromethane elimination of Br− (Equation 30.41) is followed exclusively [450]. For chlorotrinitromethane, the two pathways compete [443,450]. Fluoride ion is not a leaving group during the reduction of fluoronitroalkanes [424,441,451–454]. Thus, for example, fluorotrinitromethane is first reduced to fluorodinitromethane, which in turn may be reduced to fluoronitromethane.

–

H

R

H

+2e–

R

+ RO– –

–

R

R

(30.44)

+ NO2–

––

NO2 ––

RO

–

8. β-Substituted Nitroalkanes Studies of the electrochemical reduction of β-hydroxynitroalkanes in acidic solution may be hampered by the parallel elimination of water to the more easily reduced α,β-unsaturated nitroalkenes and in basic solution by deprotonation to the corresponding anions [455,456]. Still, it is possible to convert, for instance, methyl α-nitro-β-hydroxybutyrate to the d,l-threonine methyl ester in an overall 6F process [457,458]. Reduction of acylated β-hydroxynitroalkanes in DMF results in the formation of the alkenes by elimination of nitrite and alkoxide ions (Equation 30.44) [459].

R

R

The nitro group in 2-nitroacetophenone (a β-carbonylnitroalkane) is reduced in preference to the carbonyl group. At pH < 5, the reduction involves the enol-form according to Equation 30.45 [400,460]. Ph

N+ O

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O–

O

Ph

N+ OH

O–

O

+4e–,+4H+

Ph

H N

–H2O O

OH

(30.45)

1181

Reduction of Nitro Compounds and Related Substrates

9. α,β-Disubstituted Nitroalkanes The reduction of α,β-dibromonitroalkanes proceeds as a 2F process with the elimination of bromide to the corresponding nitroalkene (Equation 30.46) [376], whereas reduction of 1-bromo-1-nitro-2butanol leads to 1-nitro-2-butanol in a 2F process with elimination of bromide ion accompanied by the formation of 1-nitrobutene [400]. Ph

NO2

Br

Br

+2e– –2Br–

III.

Ph

CH

CH

NO2

(30.46)

NITROSO COMPOUNDS

The interest in the electrochemical reduction of nitroso compounds is partly related to the fact that they are notable intermediates in the reduction of nitro compounds as discussed earlier. Another driving force is the application of nitroso compounds as spin traps [461–466]. In this section, we will discuss only reactions in which the nitroso compounds are the starting materials. In spite of the fact that nitroso compounds are intermediates in the reduction of nitro compounds, significant differences are often observed when the reduction sequence is initiated at the nitroso level. For instance, aliphatic nitroso compounds and o-substituted nitrosobenzenes exist in solution mainly as the N-N dimers [146,462,467–472] and this has consequences for the electrochemical behavior. Also, the radical anions of nitrosobenzenes tend to dimerize (reversibly) and this of course affects the competition between the formation of hydroxylamines and azoxybenzenes. Finally, it should be recalled also that nitroso compounds are susceptible to nucleophilic attack and this is of importance when reduction is carried out under nonacidic conditions [473]. These aspects, to be discussed briefly later, are usually of no concern when the nitroso compounds result from reduction of the corresponding nitro compounds since in that case the “side reactions” do not have the time to manifest themselves, mechanistically or product-wise, before the nitroso compounds are further reduced. A summary of the electrochemistry of nitroso compounds is included in a 1994 review [473].

A. AROMATIC NITROSO COMPOUNDS Compounds in which one or more nitroso groups are attached to a benzene ring or a substituted benzene ring are included in this section. When an aromatic compound contains both a nitro and a nitroso group, the nitroso group is reduced first, usually leaving the nitro group unaffected. Such mixed compounds are therefore considered as nitroso compounds and treated in this section. 1. Formation of and Properties of the Radical Anions and Dianions The electrochemical reduction of nitrosobenzene and substituted nitrosobenzenes to the radical anions has been observed by CV in aprotic solvents such as MeCN [465,474,475], DMF [15,41,48,462,476–482], and DMSO [475] and in liquid ammonia [23]. The radical anions may, of course, be formed during CV by reduction of the monomeric nitroso compounds (Equation 30.47) but also by reductive cleavage of the dimers (Equation 30.48) [462]. It follows that the dissociation of the dimers usually is slow at the time scale of routine CV. +e– R

(R

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NO

NO)2

–e– +2e–

R

NO

2R

(30.47)

–

NO

–

(30.48)

1182

Organic Electrochemistry

The radical anions derived from nitrosobenzenes are as the rule more reactive than the corresponding nitrobenzene radical anions [15,23,478]. For example, the nitrosobenzene radical anion is a stronger base than the nitrobenzene radical anion [23], and this of course strengthens the requirements of conducting the experiments under “water-free” conditions if protonation is to be avoided. However, even under “dry” conditions, (reversible) dimerization of the radical anions, an inherent property of these species, is frequently observed [48,474–480] (see Section III.A.3). The effects of ion-pair formation follow the common rules; the stronger the ion-pair, the easier the nitroso compound is to reduce [480]. Ion-pair formation is usually accompanied by charge localization and thereby by spin localization. Thus, ion-pair formation accentuates the tendency of nitrosobenzene radical anions to dimerize. Reduction by CV of nitrosobenzene and substituted nitrosobenzenes in basic aqueous solution have in few cases resulted in the reversible formation of the radical anions [497] and the nitrosobenzene radical anion has been detected in aqueous solution by pulse-radiolysis as a short-lived species [483]. However, mostly reduction in aqueous solution proceeds to the hydroxylamine as a reversible 2e−, 2H+ process (Equation 30.8). The reduction of the radical anion to a persistent dianion requires the strict absence of residual water and has to the best of our knowledge not yet been observed at room temperature for unsubstituted nitrosobenzene: However, the reversible one-electron couple for the reduction of the radical anion has been observed in liquid ammonia [23]. Less reactive dianions are observed for nitrosobenzenes carrying an additional electron-withdrawing group such as a nitro group [41,48]. The reduction potentials have been recorded for a large number of substituted nitrosobenzenes in various solvents [83,462,465,484]. The nitrosobenzenes are easier to reduce than the corresponding nitrobenzenes under similar condition [23,41,83,133,143], and the monomers are more easily reduced than the dimers [146,462]. We are not aware of studies in which the heterogeneous electron transfer rates of nitrobenzenes and nitrosobenzenes are compared under identical conditions. However, the formation of the nitrosobenzene radical anion appears to be quasireversible at a CV sweep rate of 30 Vs−1 [474]. The electron transfer rate for a series of 4-(3-nitrosophenyl)-1,4-dihydropyridine derivatives was found to be two orders of magnitude smaller on glassy carbon than on mercury electrodes [485]. The UV/Vis [474,475,479] and ESR [148,465,466,475,478,479,486,487] spectra have been recorded for the radical anions of a number of nitroso compounds The preparative electrochemical reduction of nitroso compounds is only of practical use in rare cases owing to the limited availability of the starting materials and separate sections are for that reason not called for. Thus, the following sections will include both mechanistic and preparative aspects. 2. Routes to Phenylhydroxylamines and Anilines a. Protonation in Aprotic Solvents The initial protonation of the nitrosobenzene radical anions that ultimately leads to the formation of phenylhydroxylamines and anilines has to be fast in order to compete with the tendency of the radical anions to dimerize. Thus, it is typically observed that addition of an acid of suitable strength in increasing amounts results in a gradual change from the formation of the azoxybenzene to the formation of the hydroxylamine [23,474]. Another example that demonstrates that even subtle differences in the reaction conditions may influence the competition between dimerization and protonation is the effect of the nature of the supporting electrolyte observed for reduction of nitrosobenzene in DMF [488]. When Et4NClO 4 was used, the route to phenylhydroxylamine was found to dominate, whereas when alkalimetal perchlorates were used, the route to azoxybenzene was dominating, in the last case likely to be governed by the formation of ion-pairs [480] that localizes the spin and thereby the tendency of the radical anions to dimerize.

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Reduction of Nitro Compounds and Related Substrates

b. Protonation in Water and Other Protic Solvents The reduction of nitrosoarenes in aqueous solutions at all but the most high pH values proceeds as a 2e−, 2H+ reaction to the corresponding hydroxylamine [143,160,196,473,479,481,484,485,489,490] that during polarography and CV is observed as a reversible or quasireversible redox couple (Equation 30.8). The sequence of electron and proton transfers [160] has already been discussed in Scheme 30.6. At pH < 4, the reaction proceeds to the protonated hydroxylamine [484]. The hydroxylamine may in turn be reduced to the aniline in another 2e−, 2H+ process (see Section IV). Similarly to what was observed for the nitrobenzenes, the reduction of nitrosobenzenes substituted in the p-position with a hydroxy [206,219,484,491] or an amino group [220,221,484,492] proceeds directly to the amine in yields that may exceed 90%. 1-Nitroso-2-naphthol reacts in the same way [493,494]. The rate constants, k, for the dehydration steps (Equation 30.19) were found to be in the range 0.4−1.0 s−1 [219–221]. 3. Routes to Azoxybenzenes, Azobenzenes, Hydrazobenzenes, and benzidines The formation of the new N–N bond in the route to azoxybenzenes and the reduction products may in principle occur via (1) reduction followed by dimerization or (2) dimerization followed by reduction. In the latter case, reduction proceeds all the way to the hydrazobenzene stage owing to the low potential required to reduce the dimer, whereas electrolysis of the more easily reduced monomers may stop at the azoxy- and azo stage. The electrochemical process is accompanied by a nonelectrochemical route to azoxybenzenes initiated by the reaction between the nitrosobenzene and alkoxide or hydroxide ion [473,477,495–498], the latter being formed during electrochemical reductions even in aprotic solvents owing to deprotonation of residual water. Several mechanism suggestions for the reaction have been offered, but all the details are not yet fully understood. In addition to the azoxybenzene that results from the initial nucleophilic attack of hydroxide ion on the nitroso function [495], variable amounts of the quinone monoxime (or the anion = p-nitrosophenolate) resulting from competitive attack of hydroxide ion at the p-position are formed as well [496]. This has led to the suggestion of the stoichiometry shown in Equation 30.49 [477]. As a result of this nonelectrochemical route to azoxybenzene, the electrochemical process often consumes only 0.3–0.8F. 3Ph

NO + OH–

Ph

N+ N

Ph + –O

NO + H2O

(30.49)

O–

a. Reduction in Aprotic Solvents Azoxybenzene is detected by CV during the reduction of PhNO in MeCN [15,474,475], and it is well known that electrolysis of nitrosobenzenes in an aprotic solvent leads to the formation of the azoxybenzenes [476,479]. The yields of azoxybenzene in MeCN and DMF are typically 70–85% [477]. These results illustrate that dimerization of the electrogenerated nitrosobenzene radical anions is the default reaction; the dimer dianions are subsequently protonated by residual water (Equation 30.14) or deliberately added acids, resulting in the formation of the corresponding azoxybenzenes [23,474–481]. The route from the radical anion to the azoxybenzene product may vary from the one given by Equation 30.14, and several variants of the order of protonation and elimination of hydroxide ion have been suggested, including dimerization of protonated radical anions, depending on the proton availability. When the reaction is carried out with alkali metal salts as supporting electrolytes, ion-pairs are involved in all the reaction steps and the formation of azoxybenzene has been suggested to include the elimination of the alkali metal oxide [480]. If acids are added, the azoxybenzene may arise via the classical reaction between nitrosobenzene and phenylhydroxylamine [23]. In a special case, the reduction of nitrosophenyl-1,4-dihydropyridines, it has been suggested that the protons required for the

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formation of the azoxy compound originate from the N–H group of the parent compound in a self-protonation–type reaction [478]. The rate constants for dimerization, k dim, have been determined in a number of cases [475,478,479] the values for o- and m-nitrosotoluene radical anion being 4100 and 2700 M−1 s−1, respectively [479]. The results from a spectroelectrochemical study have indicated that attack of the radical anion in p-position of the parent nitrobenzene to produce N-(4-nitrosophenyl)-N-phenylhydroxylamine may proceed in parallel to the azoxybenzene pathway [474]. b. Reduction in Water and Other Protic Solvents As discussed earlier, the reduction of nitrosobenzenes in aqueous solution at most pH values results as the rule in the formation of phenylhydroxylamines. However, azoxybenzenes may be formed at pH > 10 but, as discussed earlier, in that case studies are hampered by the parallel nonelectrochemical reaction between nitrosobenzene and hydroxide ion to azoxybenzene. It has been suggested that the formation of azoxybenzenes under those conditions in special cases may follow the classical reaction between the nitrosobenzene and the hydroxylamine [481]. During CV of o-nitrosobenzoic acid in an aqueous 1:1 H2PO4−/HPO42− buffer, the corresponding azobenzene/hydrazobenzene reversible redox couple was observed when the scan was reversed after reduction of the dimer has taken place according to Equation 30.50 [146]. This is an example of reduction of a neutral nitrosobenzene dimer. O– O– Ar

N+ N+ Ar

+6e–,+6H+ –2H2O

Ar

H N

H N

–2e–,–2H+ Ar

+2e–,+2H+

Ar

N

N

(30.50)

Ar

4. Other Reactions of Nitrosobenzenes Examples of using nitroso compounds as spin traps include trapping of the neutral free radicals resulting from reduction of aliphatic halides [463] and pyrylium ions [464] by nitrosodurene and related nitrosobenzenes. The electrochemical reduction of nitrosobenzene in THF in the presence of a carbon acid such as fluorene or indene induces a chain reaction resulting in a mixture of the anil, the nitrone and azoxybenzene [497,499]. At low temperature, the anil is formed almost exclusively. When phenylacetylene is used as the carbon acid, azoxybenzene is the main product. Reduction of nitrosobenzene in MeCN in the presence of acetic anhydride results in the formation of the corresponding N-acetoxy-N-phenylacetamide. This reductive acylation is believed to involve the steps shown in Equation 30.51 [82] (R ═ Ph): O +e– R

NO

–e–

R

NO

–

Ac2O –AcO–

+e– R

NOHAc

R

–

NOHAc

CH3

Ac2O –AcO–

R

N

CH3

(30.51)

O O

5. Nitronitrosobenzenes The reduction of the three isomeric nitronitrosobenzenes has been investigated in DMF in the presence of various carboxylic acids [48]. Mixtures of the corresponding hydroxylamines and azoxybenzenes were obtained, in all cases resulting from reduction of the nitroso function. The azoxybenzenes were favored in general, but the stronger the acid or the higher the concentration, the higher the yield of the hydroxylamine. For 2-nitronitrosobenzene in the presence of a 10-fold excess of benzoic acid, for example, the hydroxylamine was the main product that could be obtained

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Reduction of Nitro Compounds and Related Substrates

in 80% yield. The azoxy compounds were suggested to arise via the reversible dimerization of the nitronitrosobenzene radical anions. The resulting dimer dianions were then trapped by acid. 6. Heteroaromatic and Related Nitroso Compounds Reduction of 4-nitrosopyridine-N-oxide in aqueous solution between pH 2 and pH 9 results in the formation of the azoxy derivative [500,501]. The electrochemical reduction of 6-amino-5-nitroso-1,3-dimethyluracil to 5,6-diamino-1,3dimethyluracil, an intermediate in the synthesis of caffeine and theophylline, has been investigated intensely, and it was found that the conversion may be carried out in good yield at a variety of cathodes including foamed nickel [502,503], platinum [503], Kh18N10T stainless steel [504], as well as cathodes made of carbon fibers [505]. Related to this is the tin-ion catalyzed reduction of 2,4,6-triamino-5-nitrosopyrimidine to 2,4,5,6-tetraaminopyrimidine [506], a key substrate in the synthesis of methotrexate.

B. ALIPHATIC NITROSO COMPOUNDS Studies of the electrochemical generation of the radical anions of nitrosoalkanes seem to be limited to the reduction of the tertiary 2-methyl-2-nitrosopropane [461,462,507] The radical anion has been characterized by ESR spectroscopy [465,507]. 1. Primary and Secondary Nitrosoalkanes Primary and secondary nitrosoalkanes are unstable compounds and tautomerize readily in aqueous solution to the corresponding oximes in an acid/base catalyzed reaction. As a consequence, the reduction products are those of the oximes, that is, the corresponding amines [380]. When monomer-dimer equilibrium for the nitrosoalkanes favors the dimer, reduction leads to the hydrazo compounds in a 6F process as observed, for instance, for nitrosocyclohexane (Equation 30.52) [467]. This is in analogy to what is observed for the reduction of o-nitrosobenzoic acid [146] but in contrast to the reversible 2e−, 2H+ reduction of nitroso monomers to hydroxylamine (Equation 30.8): O– NO

N+

2 H

H

H

+6e–,+6H+

N+

H N

–2H2O H

–O

H N H

(30.52)

2. Tertiary Nitrosoalkanes The 2-methyl-2-nitrosopropane monomer is reduced in unbuffered aqueous solution or in MeCN to N-tert-butylhydroxylamine [461], possibly with residual water serving as the proton donor in MeCN. The reductive acylation described earlier (Equation 30.51) may be carried out also with 2-methyl2-nitrosopropane [82] (R ═ t-Bu). 3. α-Halonitrosoalkanes The reduction of mono- and dihalogenated nitrosoalkanes proceeds analogously to the reduction of the corresponding nitroalkanes; thus 2-chloro-2-nitrosopropane gives acetoxime in 98% yield upon reduction at pH 3 (Equation 30.53) [423] and 1,1-dichloro-1-nitrosoethane gives chloroacetaldoxime in 76% yield [423].

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H3C

X

H3C

NO

+2e–,+H+

H3C NOH

–X– H3C

X = Cl, Br

(30.53)

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Organic Electrochemistry

4. Nitronitrosoalkanes (Pseudonitroles) Pseudonitroles, such as 2-nitro-2-nitrosopropane, are reduced in weak acidic aqueous solution in a 2F process, resulting in the elimination of nitrite ion and the formation of the oxime (Equation 30.54) [423]. H3C H3C

NO

+2e–,+H+

H3C NOH +

(30.54)

NO2–

H3C

NO2

IV. HyDROXyLAMINES A.

AROMATIC HYDROXYLAMINES

The electrochemical reduction of simple arylhydroxylamines in nonaqueous solvents does not seem to have received attention. However, the reduction of 2- and 4-nitrophenylhydroxylamines in DMF has been reported to proceed to the corresponding nitroanilines [41,508] in a self-protonation process reflecting the acidity of the hydroxylamine [509]. Polarographic studies of phenylhydroxylamine in aqueous solution have shown that reduction involves either the mono- or diprotonated form depending on the acidity of the solvent and that adsorption plays an important role [139,140]. The proposed reaction sequence accounting for the formation of aniline, and as side products hydrazobenzene and benzidine, is shown in Scheme 30.15 without specifying the adsorption equilibria.

B. ALIPHATIC HYDROXYLAMINES Aliphatic hydroxylamines are not easily reduced electrolytically. In acidic aqueous solution, the polarographic wave is masked by hydrogen evolution, but a reduction wave is seen in a narrow pH region around 7. Indirect reduction to amines may be accomplished by electrochemically generated Ti3+ or Fe2+ [510].

+H+ Ph

NHOH

–H+

Ph

NH2OH+

NH

+e– Ph

NH

Ph

NH–

+e–

SCHEME 30.15

NH2 +

–H2O

+H+

+H+ Ph

–H+

NH2OH

NH

+H+ Ph

Ph

NH2

Ph -H+

NH3+

H2N

Reduction pathways for phenylhydroxylamine in aqueous solution.

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NH3+

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Reduction of Nitro Compounds and Related Substrates

V. N-NITRAMINES, N-NITROSAMINES, N-NITRAMIDES, AND N-NITROSAMIDES A.

N-NITRAMINES AND N-NITROSAMINES

The reduction of primary nitramines in MeCN proceeds via a father–son-type reaction with a stoichiometry corresponding to 1F (Equation 30.55), owing to the presence of the acidic α-protons [511]. In the presence of a proton donor such as PhOH, the stoichiometry changes as expected to 2F (Equation 30.56): +2e– 2R

R

NO2

NH

+2e– R

NH

NO2

N–

NH2 + R

NO2 + NO2–

(30.56)

NH3 + NO2–

R PhOH

(30.55)

Secondary nitramines are reduced in MeCN and DMF in an irreversible two-electron process that includes cleavage of the N–N bond (Equation 30.57) [512,513]. It has been proposed that the reduction proceeds via a nitramine radical anion in which the N–N bond is considerably weakened. Protonation of the R1R2N− anion leads to the secondary amine, R1R2NH. R1 N

NO2

+2e–

R2

R1

(30.57)

N– + NO2– R2

In acidic aqueous solution, protonated N-nitramines are reduced to the unsymmetric hydrazine in a six-electron process (Equation 30.58) [514–517] with the amine as a by-product [517]. R1

+6e–,+7H+ N

R1

NO2

N

R2

NH3+ + 2H2O

(30.58)

R2

In alkaline solution, most N-nitro derivatives of primary amines are deprotonated to R–N–NO2− that cannot be reduced, with nitraminopyridines as exceptions [517], whereas N-nitramines of secondary amines are reduced in two steps. [514]. An N-nitrosamine is produced during the first reduction step in a 2F process. It is noteworthy that N-nitrosamines in basic solution are more difficult to reduce than the corresponding N-nitramines. N-Nitro-N-arylhydrazones such as 3-(N-nitro-N-arylhydrazono)pentane-2,4-diones are reduced in acidic solution in two 4F steps; first, the nitro group is reduced to the hydroxylamine state, which is followed by reductive cleavage of the C═N–N bond (Equation 30.59) [518]. Apparently, reductive cleavage of the N–N bond does not take place in the first step. The resulting 3-aminopentane-2,4dione may be further reduced to pentane-2,4-dione. NO2 Ar

N

COCH3

NHOH

+4e–,+4H+ Ar –H2O

N

COCH3

N N

COCH3

COCH3

COCH3

+4e–,+4H+ Ar –H2O

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NH

NHOH + H2N

CH COCH3

(30.59)

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Organic Electrochemistry

NH

O

N

O

+2e–

(MeCN,H2O)

+4e–

(MeCN,HOAc)

NO

N

O

SCHEME 30.16

NH2

The 2 and 4F processes in the reduction of N-nitrosamines.

Reduction of 1-nitropyrazol in acidic media results in a cleavage of the N–N bond and formation of nitrous acid; at pH > 4, a competitive reduction to the nitrosamine takes place [331]. The reduction of secondary N-nitrosamines in aprotic solvents (MeCN, DMF) in the presence of protons donors may take two different routes depending on acidity of the solvent system. In the presence of weak acids (water, PhOH), a two-electron process including N–N cleavage and the formation of the amines is observed [519] (Scheme 30.16), similar to what has been observed in aqueous base [520], whereas the preferred route in the presence of stronger acids (AcOH, PhCOOH) leads to the formation of 1,1-dialkylhydrazines in a four-electron process [519]. This effect of acidity has been studied in detail in aqueous solution, and it was found that the 2F process leading to the secondary amine is accompanied by the formation of N2O (Equation 30.60), whereas the 4F process leading to the hydrazines involves the protonated N-nitrosamine (Equation 30.61) [521–524]. The latter reaction is a convenient synthetic route to 1,1-dialkylhydrazines [524–526]. R1

R1

+4e– N

2

NO + 3H2O

2

NH + N2O + 4OH–

(30.60)

NH3+ + H2O

(30.61)

R2

R2 R1 N

NOH+ + 4H+

R2

+4e–

R1 N R2

The nitrosation of secondary amines in acidic solution is a reversible reaction and, accordingly, N-nitrosamines may undergo hydrolysis at low pH. This is especially important for diaryl N-nitrosamines, but also significant for aryl alkyl N-nitrosamines. The NO+ liberated during hydrolysis may react with the reduction product from the N-nitrosamine, the unsymmetrical hydrazine, with the formation of the secondary amine and N2O usually resulting under basic conditions [527] (Equation 30.62). R1

R1 N

NH2

+ NO

+

R2

B.

NH

+ N2O + H+

(30.62)

R2

N-NITRAMIDES AND N-NITROSAMIDES

The electrochemical reduction of N-nitramides and N-nitrosamides proceeds in essentially the same manner as described earlier for the N-nitramines and N-nitrosamines. In aprotic media such as DMF, the primarily formed radical anion of N-nitrosourea is protonated by the substrate giving rise to a cascade of follow-up reactions including further reduction of the protonated radical anion [528] (Equation 30.63).

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1189

Reduction of Nitro Compounds and Related Substrates R

+e–

O

N

H2N

O

N

(30.63)

O

O R

R OH

N

H2N

N

O–

N

H2N

N

O

N

H2N R

R

O

+

N

NH–

N N

O

O

N-nitrosamides are reduced in acidic solution to the acylated hydrazines, which again may be hydrolyzed to the corresponding alkylhydrazine. This has resulted in the following sequence (Equation 30.64) as a convenient route from primary amines to the corresponding hydrazines [529]. R

NH2

CH3COOH >120°C,–H2O

R

NHCOCH3

NaNO2,HCl

R

–H2O

NCOCH3 NO

(30.64) H+,H2O

+4e–,4H+ –H2O

R

R

NCOCH3

NHNH2

+ CH3COOH (Recycled)

NH2

Similarly, nitro and nitroso derivatives of urea and guanidine have been reduced to semicarbazides [530,531] and aminoguanidines [516,532–535] in acidic solution.

VI.

NITRIC ACID ESTERS

Studies of the electrochemical reduction of nitric acid esters (organic nitrates) in aprotic solvents are rare [353,536]. The reduction of 2-substituted 1-oxo-2,3-dihydro-1H-inden-2-yl nitrate in DMF proceeds as an overall 2F process, resulting in concurrent C–O and N–O cleavage [536] (Scheme 30.17). In aqueous solution, organic nitrates are reduced in a pH-independent 2F process to nitrite ion and, after protonation, the corresponding alcohol [537–539] (Equation 30.65). R

ONO2

+2e– –

R

O–

H+ R

(30.65)

OH

–NO2

OH 38–50%

+2e– ONO2 –NO2– ,+H+ R O

R O

+2e– H

–NO3– ,+H+

16–18% R O

SCHEME 30.17 Competition between C-O and N-O cleavage during the reduction of a 2-substituted 1-oxo2,3-dihydro-1H-inden-2-yl nitrate

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Organic Electrochemistry

It has been observed that the reduction of propylene glycol 1,2-dinitrate proceeds readily at cathodes made from silver, gold, copper, and mercury, but not at platinum, iridium, nickel, tungsten, and molybdenum cathodes. It was proposed that the process includes first a one-electron reduction of the organic nitrate to RO − and NO2, the latter in an adsorbed state, and then reduction of (NO2)ads to free NO2− in a second one-electron step [539].

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457. Novikov, V.T.; Ratnikova, L.A.; Avrutskaya, I.A.; Fioshin, M.Y.; Belikov, V.M.; Babievskii, K.K. Elektrokhimiya 1976, 12, 1066. 458. Novikov, V.T.; Avrutskaya, I.A.; Fioshin, M.Y.; Belikov, V.M.; Babievskii, K.K. Elektrokhim. 1976, 12, 1061. 459. Petsom, A.; Lund, H. Acta Chem. Scand. 1980, B34, 6140. 460. Deswarte, S. Bull. Soc. Chim. Fr. 1969, 534. 461. Mclntire, G.L.; Blount, H.N.; Stronks, H.J.; Shetty, R.V.; Janzen, E.G. J. Phys. Chem. 1980, 84, 916. 462. Gronchi, G.; Courbis, P.; Tordo, P.; Mousset, G.; Simonet, J. J. Phys. Chem. 1983, 87, 1343. 463. Martigny, P.; Mabon, G.; Simonet, J.; Mousset, G. J. Electroanal. Chem. 1981, 121, 349. 464. Klima, J.; Volke,, J.; Urban, J. Electrochim. Acta 1991, 36, 73. 465. Gronchi, G.; Tordo, P. Res. Chem. Intermed. 1993, 19, 733. 466. Cerri, V.; Frejaville, C.; Vila, F.; Allouche, A.; Gronchi, G.; Tordo, P. J. Org. Chem. 1989, 54, 1447. 467. Schindler, R.; Lüttke, W.; Holleck, L. Chem. Ber. 1957, 90, 157. 468. Fletcher, D.A.; Gowenlock, B.G.; Orrell, K.G. J. Chem. Soc. Perkin Trans. 1997, 2, 2201. 469. Fletcher, D.A.; Gowenlock, B.G.; Orrell, K.G. J. Chem. Soc. Perkin Trans. 1998, 2, 797. 470. Fletcher, D.A., Gowenlock, B.G.; Orell, K.G.; Apperley, D.C.; Hursthouse, M.B.; Malik, K.M.A. J. Chem. Res. (S) 1999, 202. 471. Stowell, J.C. J. Org. Chem. 1971, 36, 3055. 472. Holmes, R.R. J. Org. Chem. 1964, 29, 3076. 473. Zuman, P.; Shah, B. Chem. Rev. 1994, 94, 1621. 474. Steudel, E.; Posdorfer, J.; Schindler, R.N. Electrochim. Acta 1995, 40, 1587. 475. Nuñez-Vergara, L.J.; Squella, J.A., Olea-Azar, C.; Bollo, S.; Navarrete-Encina, P.A.; Sturm, J.C. Electrochim. Acta 2000, 45, 3555. 476. Kemula, W.; Sioda, R. J. Electroanal. Chem. 1963, 6, 183. 477. Asirvatham, M.R.; Hawley, M.D. J. Electroanal. Chem. 1974, 57, 179. 478. Nuñez-Vergara, L.J.; Santander, P.; Navarrete-Encina, P.A.; Valenzuela, J.; Sturm, J.C.; Squella, J.A. J. Electrochem. Soc. 2006, 153, E144. 479. Nuñez-Vergara, L.J.; Bontá, M.; Sturm, J.C.; Navarrete, P.A.; Bollo, S.; Squella, J.A. J. Electroanal. Chem. 2001, 506, 48. 480. Lipsztajn, M.; Krygowski, T.M.; Laren, E.; Galus, Z. J. Electroanal. Chem. 1974, 57, 339. 481. Núñez-Vergara, L.J.; Bollo, S.; Fuentealba, J.; Sturm, J.C.; Squella, J.A. Pharmaceutical Res. 2002, 19, 522. 482. Dickerson, R.L.; Rogers, J.W. Anal. Chim. Acta 1974, 71, 433. 483. Asmus, K.D.; Beck, G.; Henglein, A.; Wigger, A. Ber. Bunsenges. Phys. Chem. 1966, 70, 869. 484. Holleck, L.; Schindler, R. Z. Electrochem. 1956, 60, 1138. 485. Bollo, S.; Finger, S.; Sturm, J.C.; Núñez-Vergara, L.J.; Squella, J.A. Electrochim. Acta 2007, 52, 4892. 486. Geels, E.J.; Konaka, R.; Russell, G.A. Chem. Commun. (Lond.) 1965, 13. 487. Ayscough, P.B.; Sargent, F.P.; Wilson, R. J. Chem. Soc. B 1966, 903. 488. Lipsztajn, M.; Krygowski, T.M.; Galus, Z. J. Electroanal. Chem. 1977, 81, 347. 489. Chuang, L.; Fried, I.; Elving, P.J. Anal. Chem. 1964, 36, 2426. 490. Smith, J.W.; Waller, J.G. Trans. Faraday Soc. 1950, 46, 290. 491. Muralidharan, S.; Chellammal, S.; Anantharaman, P.N. Bull. Electrochem. 1991, 7, 222. 492. Polat, K.; Aksu, M.L.; Pekel, A.T. J. Appl. Electrochem. 2000, 30, 733. 493. Vasudevan, D.; Anantharaman, P.N. J. Appl. Electrochem. 1994, 24, 559. 494. Khatri, O.M.P.; Sharma, R.; Kumbhat, S. Bull. Electrochem. 2003, 19, 477. 495. Zuman, P.; Fijalek, Z. J. Org. Chem. 1991, 56, 5486. 496. Lund, H.; Skov, K.; Pedersen, S.U.; Lund, T.; Daasbjerg, K. Coll. Czech. Chem. Commun. 2000, 65, 829. 497. Mugnier, Y.; Gard, J.C.; Huang, Y.; Couture, Y.; Lasia, A.S.; Lessard, J. J. Org. Chem. 1993, 58, 5239. 498. Hutton, J.; Waters, W.A. J. Chem. Soc. B 1968, 191. 499. Williot, F.; Bernard, M.; Lucas, D.; Mugnier, Y.; Lessard, J. Can. J. Chem. 1999, 77, 1648. 500. Roffia, S.; Raggi, M.A.; Ciano, M. J. Electroanal. Chem. 1975, 62, 403. 501. Roffia, S.; Raggi, M.A. J. Electroanal. Chem. 1976, 67, 11. 502. Hu, X.E., Yang, H.W.; Wang, X.J.; Bai, R.S. J. Appl. Electrochem. 2002, 32, 321. 503. Chen, R.; Zheng, X.; Hu, X. Res. Chem. Intermed. 2012, 38, 1119. 504. Konarev, A.A. Russ. J. Electrochem. 2007, 43, 1206. 505. Chen, R. Res. Chem. Intermed. 2012, 38, 2111. 506. Arkhipova, T.A., Avrutskaya, I.A. Russ. J. Electrochem. 1996, 32, 114. 507. Sosonkin, I.M.; Belevskii, V.N.; Strogov, G.N.; Domarev, A.N.; Yarkov, S.P. Zhur. Org. Khim. 1982, 18, 1504.

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508. Syroeshkin, M.A.; Mendkovich, A.S.; Mikhalchenko, L.V.; Rusakov, A.I.; Gul’tyai, V.P. Mendeleev Commun. 2009, 19, 258. 509. Bordwell, F.G.; Liu, W.-Z. J. Am. Chem. Soc. 1996, 118, 8777. 510. Feroci, G.; Lund, H. Acta Chem. Scand. 1976, B30, 651. 511. (a) Semakhina, N.I.; Podkovyrina, T.A.; Supyrev, A.V.; Lyzhina, L.I.; Kargin, Yu.M. Zhur. Obshchei Khim. 1982, 52, 2316; (b) Petrosyan, V.A.; Frolovsky, V.A.; Sadilenko, D.A. Russ. Chem. Bull. 1999, 48, 83. 512. Yanilkin, V.V.; Berdnikov, E.A.; Buzykin, B.I. Russ. J. Electrochem. 2000, 36, 144. 513. (a) Kargin, Yu.M.; Kondranina, V.Z.; Gafarov, A.N.; Ivshin, V.P.; Podkovyrina, T.A.; Semakhina, N.I.; Kazakova, A.A.; Yanilkin, V.V.; Koloskova, T.N.; Vakhrusheva, E.M. Zhur. Obshchei Khim. 1977, 47, 666; (b) Kargin, Yu.M.; Latypova, V.Z.; Supyrev, A.V. Zhur. Obshchei Khim. 1982, 52, 2623; (c) Kargin, Yu.M.; Marchenko, G.A.; Latypova, V.Z.; Punegova, L.N.; Supyrev, A.V.; Bogoveeva, G.A.; Egorova, L.S.; Stepanov, G.S. Zhur. Org. Khim. 1986, 22, 45; (d) Semakhina, N.I.; Podkovyrina, T.A.; Toktaulova, L.O.; Kargin, Yu.M. Zhur. Obshchei Khim. 1986, 56, 2764; (e) Semakhina, N.I.; Podkovyrina, T.A., Shabalin, A.F.; Kargin, Yu.M. Zhur. Obshchei Khim. 1984, 54, 2103. 514. Laviron, E.; Fournari, P. Bull. Soc. Chim. Fr. 1966, 518. 515. Laviron, E.; Fournari, P.; Greusard, M. Bull. Soc. Chim. Fr. 1967, 1255. 516. Laviron, E.; Fournari, P.; Refalo, G. Bull. Soc. Chim. Fr. 1969, 1024. 517. Lund, H.; Sharma, S.K. Acta Chem. Scand. 1972, 26, 2329. 518. Jain, R.; Agarwal, D.D.; Shrivastava, R.K. J. Chem. Soc. Perkin Trans. 1990, 2, 1353. 519. (a) Kargin, Yu.M.; Latypova, V.Z.; Supyrev, A.V.; Kucherova, N.L. Zhur. Obshchei Khim. 1982, 52, 338; (b) Kargin, Yu.M.; Latypova, V.Z.; Supyrev, A.V.; Zhuikov, V.V. Zhur. Obshchei Khim. 1984, 54, 1695. 520. Hlophe, M. Int. J. Electrochem. Sci. 2012, 7, 5927. 521. Lund, H. Acta Chem. Scand. 1957, 11, 990. 522. Holleck, L.; Schindler, R. Z. Elektrochem. 1958, 62, 942. 523. (a) Pulidori, F.; Borghesani, G.; Bighi, C.; Pedriali, R. J. Electroanal. Chem. 1970, 27, 385; (b) Borghesani, G.; Pulidori, F.; Pedriali, R.; Bighi, C. J. Electroanal. Chem. 1971, 32, 303. 524. (a) Nikulin, V.N.; Klochkova, V.N. Elektrokhimiya 1972, 8, 499; (b) Subbiah, P.; Noel, M.; Chidambaram, S.; Udupa, K.S. Bull. Electrochem. 1987, 3, 181. 525. Iversen, P.E. Acta Chem. Scand. 1971, 25, 2337. 526. (a) Iversen, P.E. Chem. Ber. 1972, 105, 358; (b) Pachori, R.; Mishra, S.C.; Mishra, R.A. J. Electrochem. Soc. India 1985, 34, 99; (c) Nedungadi, P.A.K.; Gupta, A.; Mukherji, S.K.; Zutshi, K. J. Electrochem. Soc. India 1986, 35, 203. 527. (a) Jacob, G.; Moinet, C.; Tallec, A. Electrochim. Acta 1982, 27, 1417; (b) Jacob, G.; Moinet, C.; Tallec, A. Electrochim. Acta 1983, 28, 635. 528. Escot, M.T.; Martre, A.M.; Pouillen, P.; Martinet, P. Bull. Soc. Chim. Fr. 1986, 548. 529. Moore Jr., M.P. US Pat 3,267,011, 1966. 530. Won, M.S., Kim, J.C.; Shim, Y.B. J. Korean Chem. Soc. 1991, 35, 707. 531. Won, M.S., Kim, J.C.; Shim, Y.B. Bull. Korean Chem. Soc. 1992, 13, 214. 532. Yamashita, M., Sugino, K. J. Electrochem. Soc. 1957, 104, 100. 533. Pathy, M.S.V. Electrochem. Technol. 1965, 3, 94. 534. Spreter, V.C.; Briner, E. Helv. Chim. Acta 1949, 32, 215. 535. (a) Won, M.S.; Kim, J.C.; Shim, Y.B. J. Korean Chem. Soc. 1991, 35, 707; (b) Won, M.S.; Kim, J.C.; Jeong, E.D.; Shim, Y.B. J. Korean Chem. Soc. 1995, 39, 842. 536. Orliac-Le Moing, A.; Delaunay, J.; Simonet, J. Electrochim. Acta 1987, 32, 1769. 537. Kaufman, F.; Cook, H.J.; Davis, S.M. J. Am. Chem. Soc. 1952, 74, 4997. 538. Whitnack, G.C.; Nielsen, J.M.; Gantz, E.St.C. J. Am. Chem. Soc. 1954, 76, 4711. 539. (a) Miles, M.H.; Fine, D.A. J. Electroanal. Chem. 1981, 127, 143; (b) Fine, D.A.; Miles, M.H. Anal. Chim. Acta 1983, 153, 141.

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31

Reduction of Aldehydes, Ketones, and Azomethines Jiří Ludvík

CONTENTS I. Carbonyl Compounds ............................................................................................................. 1202 A. Introduction .................................................................................................................... 1202 1. General Mechanistic Considerations ....................................................................... 1203 2. Structure–Potential Correlations ............................................................................. 1206 B. Reduction of Individual Carbonyl Compounds to Alcohols and/or Pinacols ................ 1207 1. Aliphatic Carbonyl Compounds .............................................................................. 1207 2. Aromatic Carbonyl Compounds .............................................................................. 1208 C. Reduction of Carbonyl Compounds Influenced by α-Substitution................................. 1209 1. α,β-Unsaturated Carbonyl Compounds ................................................................... 1210 2. α-Substituted Carbonyls .......................................................................................... 1214 D. Reduction of Carbonyls with a Remote Unsaturation–Intramolecular Coupling .......... 1215 1. Remote Double Bond or Aromatics ........................................................................ 1215 2. Remote Unsaturated Groups ................................................................................... 1216 E. Reduction of Carbonyls in the Presence of Another Reactant–Intermolecular Coupling .... 1217 F. Reduction of Dicarbonyl Compounds ............................................................................ 1221 1. 1,2-, 1,3-, and 1,4-Dicarbonyls on an Aromatic Ring ............................................. 1221 2. 1,2-Dicarbonyls........................................................................................................ 1222 3. 1,3-Dicarbonyls and Remote Diones ....................................................................... 1223 4. β-Keto Esters ........................................................................................................... 1224 G. Reduction Mechanisms in Analytical Applications ....................................................... 1226 II. Azomethine Compounds ........................................................................................................ 1228 A. General Mechanistic Considerations .............................................................................. 1228 B. Derivatives of Ammonia (Imines, Schiff Bases, Iminium Cations) .............................. 1229 1. Protic Media ............................................................................................................ 1229 2. Aprotic Media .......................................................................................................... 1230 C. Derivatives of Hydroxylamine (Oximes, O-Alkylated Oximes, N-alkylated Oximes–Nitrones)....................................................................................... 1233 1. Protic Media ............................................................................................................ 1233 2. Aprotic Media .......................................................................................................... 1236 3. Mono- and Dioximes of α-Diketones ...................................................................... 1237 D. Derivatives of Hydrazine (Hydrazones) ......................................................................... 1238 1. Protic Media ............................................................................................................ 1238 2. Aprotic Media .......................................................................................................... 1239 E. Azines (Cyclic, Acyclic) .................................................................................................1240 1. Protic and Aprotic Media ........................................................................................ 1240 2. Hydrazonates ........................................................................................................... 1242 References .................................................................................................................................... 1242

1201

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Organic Electrochemistry

I. CARbONyL COMPOUNDS A.

INTRODUCTION

Carbonyl compounds (aldehydes, ketones, quinones, and their nitrogen derivatives—oximes, imines, azines, and hydrazones) are broadly electrochemically investigated and thus, hundreds of contributions and several review articles and chapters were published during the last five decades (e.g., References 1–4). In this chapter, besides some fundamental information, the stress was given to more recent results (particularly from the last 15 years) to follow the development in the field. For classical original papers, please consult the respective contributions by various authors in previous editions of this monograph [5–9]. For the carbonyl group—due to its “lack of electrons” (it belongs to the electron-withdrawing groups [EWG])—reduction is its typical electrochemical process. The (electro)chemical behavior of carbonyl compounds depends strongly on the structure of the rest of the molecule (aliphatic, olefinic, aromatic, α- or β- or ω-substitution, planarity—electron delocalization, etc.) and on experimental conditions (protic/aprotic medium, pH, electrode material, presence of other reactant, etc.). Therefore, reduction potentials of carbonyl derivatives vary in the span of more than 2 V: for example, p-benzoquinone (−0.4 V) and acetone (−2.5 V). When only aldehydes or ketones are reduced, the pattern is in principle similar and simple. The first electron transfer generates anion radical species. This reaction is followed either by another one-electron reduction accompanied by protonation (a two-electron heterogeneous process) leading to an alcohol or by coupling (a one-electron process followed by a homogeneous reaction) yielding a pinacol [10–12] (Scheme 31.1). The order and potentials (energetics) of individual electron and proton transfers, however, differ substantially being influenced by the structure and aromaticity of the molecule bearing the carbonyl group and by experimental conditions. As a consequence, various mechanisms are involved, various intermediates are participating, and thus, different proportions of alcohol/pinacol products result. General rules: • Aldehydes are reduced at less negative potentials than analogous ketones. • Reduction of aromatic species proceeds always at less negative potentials than reduction of aliphatic ones. • Reduction in protic media (e.g., in water) occurs at less negative potentials than in aprotic solvents due to (antecedent and/or follow-up) protonation reactions. • Nitrogen derivatives of carbonyls (imines, oximes, hydrazones) and sulfur analogs (thiocarbonyls) are reduced more easily than the “true” carbonyls. Since the carbonyl group involves π-electrons, its conjugation with another, possibly present unsaturated grouping in α-position (or in para-position at an aryl ring) generates a more delocalized and

R1

R2

R1

O

O

SCHEME 31.1

R2

R2 OH OH

+2e–, +2H+

+ R1

OH

+2e–, +2H+

O

R1 R1

R2

R1 R2

R2

Two-electron vs. one-electron reduction mechanism of carbonyls.

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Reduction of Aldehydes, Ketones, and Azomethines

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more easily reducible system with more stable intermediates. The original (electro)chemical properties of the carbonyl are then changed and influenced by the rest of the molecule. In the presence of another unsaturated group (alkene, aromatics, imine, oxime, nitrile, etc.) either suitably remote in the same molecule or as a reaction partner dissolved in the solution, coupling reactions can occur being of high importance in electrosynthesis. In protic media, for example, methanol/tetrahydrofuran (MeOH/THF) or isopropylalcohol/dimethylformamide (i-PrOH/DMF), the primary carbonyl radical anion is protonated and a σ-radical located at the carbon atom is formed. This radical species attacks the unsaturated center and a new C–C bond is formed after the second electron transfer and necessary protonation. Such a radical addition can proceed either as an intramolecular cyclization or intermolecularly under formation of an asymmetric coupling product. The reduction of the carbonyl group to the alcohol is always the competing reaction. Earlier in this section, only primary reduction of the carbonyl moiety was considered. Taking into account generally the electrochemical reduction of carbonyl compounds, we can meet another possibility: the carbonyl group itself is not reduced but it is acting as an “activator,” which means its presence causes a more easy (primary) reduction of another reducible center of the same molecule. This case may occur as a consequence of a very strong intramolecular electron interaction discussed earlier when due to electron delocalization the displacement of the LUMO is changed (e.g., chalcones or benzalacetones) [13–15]. Finally, the carbonyl group represents not only an EWG but also a center of possible tautomerism or a function promoting planarity and extending delocalization. In addition to this, due to the polarity of the CO bond, formation of hydrogen bonds is possible. These additional properties must be also taken into account when discussing the electroreduction mechanism. 1. general Mechanistic Considerations From the mechanistic point of view, the fundamental (electro)chemical properties of the C=O bond itself are given by the strong polarization of the carbon–oxygen double bond with a partly positive charge at the carbon. In the presence of protons (particularly in aqueous or mixed solutions at acidic, neutral and slightly basic pH), an antecedent pre-protonation of the partly negatively charged oxygen [16] takes place before the electron transfer. In addition to this, in the presence of nucleophiles (including water) a nucleophilic addition to the C=O double bond (e.g., hydration, addition of amines, thioles, alcohols) can occur parallel with protonation and reduction [17] Since these antecedent reactions represent equilibria with various reaction rates of their restoration, the apparent kinetics of reduction of carbonyls may be very complicated and crucially depending on the reaction conditions (pH, temperature, concentration of nucleophiles, ratio of reacting components, etc.). A simplified general scheme of possible electroreduction pathways of the carbonyl group itself is presented in Scheme 31.2. In aprotic media, a pre-protonation does not take place. The lack of electrons on the carbon atom of the parent carbonyl A causes its (mostly reversible) one-electron reduction at the potential E3, yielding a radical anion B with the unpaired electron localized on the carbon. The stability of this radical intermediate depends on the delocalization possibilities: Whereas the radical anion of a saturated carbonyl molecule is highly unstable and thus reactive, the presence of an aromatic ring or olefinic system in α-position increases the stability substantially. Primary radical anions of some aromatic carbonyls then can be detected by cyclic voltammetry (CV) or by in situ EPR spectroelectrochemical experiments. At more negative potentials (E4) formation of a dianion C is possible (pathway A–B–C in Scheme 31.2). Eventual protonation of already reduced species can be caused either by traces of moisture or by proton abstraction from the solvent. The typical sequence of steps is EECC (e−, e−, H+, H+). In protic solutions, the situation is determined mainly by acid–base equilibria, which means by protonation before, during, and/or after the electron transfers. In addition to this, at a pH more basic than

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Organic Electrochemistry 2–

C O

+H+

+e– E4

B

F –

O– +e– E3 A

(+) (–) O

OH +H+

+H+

+e–

G E2

D

+H+

OH+

HC

+e–

OH

OH

E1 E

+OH– + H–Nu (H–OH; H–NHR; H–OR;...)

+e–, +H+ (x 2)

Dim

J O–

+e–

OH

K

OH Nu

H

OH

OH

Nu = NHR N-R

–H2O

+e–

L

+2e–, +2H+ HC

E5

NH–R M

SCHEME 31.2 Possible pathways of carbonyl group electroreduction. ilim

i2 i1 + i2

i3 + i2

i1 pH E1/2

E3

E2

E1

pH

FIgURE 31.1 Typical i–pH and E1/2(red)–pH plots and the link between them for cathodic two-electron reduction of carbonyl compounds in buffered aqueous solutions.

pKa by three to four units, the protonation rate becomes slow and the electroreduction process appears to be kinetically controlled. Therefore, all electrochemical steps are apparently irreversible and the pH of the solution is a very important factor for the chemoselectivity of the electroreduction mechanism. The following behavior is typical particularly for aryl carbonyls (see Scheme 31.2 and Figure 31.1): 1. In the acidic and neutral pH region, a partly negative charge on oxygen allows its antecedent protonation (pre-protonation), resulting in a positively charged molecule (D). Its respective “true” pKa is usually around 2–3; however, due to the fast acidobasic equilibrium, the reduction of carbonyls occurs solely via pre-protonated species (D–E) up to pH 6–8.

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Species D is reduced much more positively (at E1) than the parent compound (at E3). The neutral radicalic intermediate (E) can be either reduced at more negative potential E2 to an anion (F), which is protonated to alcohol (G), or it undergoes a dimerization yielding a pinacol (H). The latter pathway is favored under these conditions due to the hydroxycarbinyl radical (E) stability. In the first case, two one-electron reduction steps at E1 and E2 are observed—pathway A–D–E–F–G with the sequence CEEC (H+, e−, e−, H+), pinacolization is a one-electron process A–D–E–H, hence, CEdim (H+, e−, dim). The reduction of pre-protonated species D is always pH-dependent, where the potential (E1) is shifted to more negative values with increasing pH due to the decreasing rate of pre-protonation. This pH dependence can be used as an indication of the pre-protonation pathway. On the other hand, reduction of the neutral protonated radical E is pH-independent and proceeds at the constant potential E2. Therefore at certain (still slightly acidic) pH, these two dependencies have to cross and at pH values where E1 becomes more negative than E2, the two originally one-electron waves merge to only one irreversible two-electron reduction process (see Reference 4). With further increasing pH, the rate of pre-protonation becomes the rate-determining step (RDS) for the entire reduction. The two-electron wave (peak) is no more controlled by transport but kinetically and its current decreases to zero. Simultaneously, the process E1 + E2 is gradually replaced by a more negative reduction process (E3) corresponding to the reduction of unprotonated parent molecule (A–B). 2. In medium–basic pH (or in slightly protic organic solvents), where pre-protonation does not proceed, the parent compound is reduced at E3 (A–B) and the intermediate B is immediately protonated. The protonation (B–E) represents another acidobasic system with pKa in the range 11–12 for aliphatic carbonyls and 8–10 for aromatic ones. Since generally the protonated radical anion (E) is more easily reducible than the parent compound (A), the potential E3 is always more negative than E2 and a “classical” ECEC (e−, H+, e−, H+) mechanism proceeds and only one irreversible, pH-independent two-electron reduction under formation of alcohol is observed (pathway A–B–E–F–G). 3. Strongly basic solutions can be considered as similar to aprotic conditions since the rate of protonation B–E is very slow. Therefore, the reduction starts at E3 and at more negative potentials (E4) a dianion may be formed (A–B–C), resulting finally in the corresponding alcohol and following the EECC (e−, e−, H+, H+) pathway. In very alkaline solutions, addition of hydroxyl anion (OH−) to the carbonyl carbon atom of the parent compound A may occur, yielding a very difficultly reducible adduct J and due to that the total reduction current decreases. The typical i–pH and E1/2(red)–pH plots and the link between them are depicted in Figure 31.1. In the presence of a nucleophile (water, alcohols, ammonia or primary amine, thiol, etc.), particularly when the carbonyl is further activated by a neighbor EWG making the carbon even more electron deficient, an anteceding covalent nucleophilic addition to the C=O double bond occurs, resulting in a gem-diol (with water), hemiacetal or hemiketal (with alcohol), carbinolamine (with ammonia or amine), or similar structures (K). These adducts are generally nonreducible in the available potential range, but they should be considered as “masked” carbonyls (Scheme 31.3). In the case of water, the hydrate is in equilibrium with the parent molecule. Hence, the bulk concentration of the parent molecule is given by the equilibrium constant and thus reduction of the parent carbonyl depends on the rate of dehydration. If the dehydration is slow, the reduction itself becomes controlled by this antecedent kinetics. A novel approach to the analysis of the dehydration kinetics of several carbonyl compounds allowed obtaining correct rate constants under steady-state conditions, independently of the degree of the diffusion contribution [18]. In the case of primary amines, the carbinolamine K undergoes dehydration yielding an imine L, which is reduced more easily than the parent carbonyl (E3 is more negative than E5), giving rise to

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Organic Electrochemistry

O

H2O R1

R2

O

H2N–R3

gem-diol R2

R3–OH

O

HO OH

HO

R2

R1

O

R3

R3–OH

R3

O O

R1

HO

H N

R3

R1 R2 carbinolamine

R3 + H2O

R2

R2

R1

hemiketal (hemiacetal if R1

SCHEME 31.3

R2

R1 H)

R1

ketal (acetal if R1

H)

Various covalent nucleophilic additions to the C=O double bond.

a secondary amine (pathway A–K–L–M). In the case of alcohols, hemiacetals or hemiketals may either dissociate back to carbonyl and alcohol, or, under excess of alcohol (in acidic solution) acetals or ketals are formed. In aldehydes due to the polarity of the C=O bond, the aldehydic hydrogen is acidic and can act as a proton donor toward strong nucleophiles (like primary radical anions). This reaction can be important in aprotic media where autoprotonation may occur. Carbonyl derivatives bearing a hydrogen atom in α-position may undergo also keto-enol tautomerism [19–21]. It is worth mentioning that both forms may be electrochemically reducible. 2. Structure–Potential Correlations When studying electrochemical reduction or oxidation of a carbonyl compound bearing various substituents and/or additional electroactive groups (often as an electrochemical characterization of newly synthesized molecules), the carbonyl group may, but also may not, be reduced as the first. In addition to this, within a homologous series of substitution or structural derivatives of the parent carbonyl compound, various mechanisms may occur. In the case of a series of relatively complicated organic substances with several possible redox centers, using analysis of structure–potential correlations and substituent effects, it is possible to localize at least the first oxidation and the first reduction center of the molecule, to estimate the preferentially electrolyzed groups, to predict the redox potentials, and to interpret possible redox mechanisms, including relative rates and equilibrium constants. Among the relationships correlating measured potentials with nature of substituents, the Hammett approach is the most widely applied [22]; in electrochemistry, it is known as linear free energy relationship (LFER) [23]. Application of the LFER treatment for reduction of carbonyl derivatives is the first choice. In a series of compounds with systematically changed substitution, the reduction potentials should follow a linear plot against the substituent constants σ (sigma) where EWG (with positive σ-values) facilitate the reduction and electron-donating groups (negative σ-values) shift the potential to more negative values. All derivatives whose reduction potentials fit the linear relationship are reduced analogously, according to the same mechanism. The eventual anomalous values of reduction potentials nonfitting the straight line indicate, however, a different reduction center, different shape of LUMO, and thus different reactivity and properties worthy to be studied more deeply [24]. The extent of intramolecular electronic communication between the substituent and the reduction center is reflected in the slope of the Ered versus σ relationship, so-called reaction constant, ρ (rho) value. The experimentally acquired ρ-values and their treatment can be used for localization of redox centers for the determination of the susceptibility of individual positions toward reduction and for the evaluation of electron distribution within the molecule [25].

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1207

The LFER treatment is important particularly in cases when besides the studied carbonyl compounds the molecule involves double- or triple bonds and/or another reducible grouping (including carbonyl itself). In the case of reduction of mono- and disubstituted phenyl styryl ketones (chalcones), multiple substituent effects were analyzed using the Hammett equation and an interactive Gibbs energy relationship. A cross-interaction was described and discussed in terms of redox reactivity [14]. Seven diaryl ketones were reduced in eight aprotic solvents, and the measured potentials were treated by the Hammett approach. The obtained reaction constants correlate with the solvent acidity/basicity expressed by the acceptor/donor numbers pointing out the important role of the solvation of participating species [26]. Recently, a quantitative structure–electrochemistry relationship (QSER) approach has been used for prediction of reduction potentials in a series of 73 aldehydes and ketones. This most advanced study of relationship between the structure and reduction potentials involves multiple liner regression (MLR), partial least square (PLS), artificial neural network (ANN), and wavelet neural network (WNN) modeling methods. The mentioned nonlinear methods exhibited better predictive power than the linear ones [27].

B. REDUCTION OF INDIVIDUAL CARBONYL COMPOUNDS TO ALCOHOLS AND/OR PINACOLS First, let us consider the most simple case of carbonyl compounds: besides the carbonyl group, neither the molecule nor the electrolyzed solution involves any other reactive function. For this situation, the standard pathways are depicted in Scheme 31.1 and a mixture of alcohol (reduction product G) and pinacol (hydrodimerization product H) is obtained. Their proportion depends 1. On the type of carbonyl compounds being reduced: Aromatic derivatives are able to form a more extended delocalized system, which makes the reduction easier, stabilizes radical intermediates, and thus promotes their various coupling reactions, including pinacol formation. On the other hand, (di)alkyl carbonyl compounds are reduced at more negative potentials, their intermediates are unstable and undergo fast protonation and further reduction at the electrode, giving rise preferentially to a two-electron product, the alcohol. 2. On the solvent and pH: Protonation during reduction favors alcohol formation. On the other hand, lack of protons stabilizing the primary radical anion B (in aprotic media) facilitates pinacolization and the same effect has an antecedent protonation in strongly acidic solutions making E1 less negative than E2. 1. Aliphatic Carbonyl Compounds Generally, the main reduction product of aliphatic carbonyl substrates is the corresponding alcohol. Only in acidic media (pH 2–5) when processes E1 and E2 are still separated, the formation of pinacol is more preferred (pathway A–D–E–H; see Section I.A.1). A detailed overview of classical electrosynthetic reactions of nonolefinic aldehydes and ketones yielding alcohols is presented in Reference 6. Sugars should be also considered as saturated aliphatic aldehydes and electrochemically reduced after conversion from the hemiacetal to the acyclic form. The most important reaction is reduction of glucose to sorbitol, where a corresponding pinacol is the by-product. The reduction is either direct at a lead cathode [28] or indirect electrocatalytic hydrogenation at the Raney nickel cathode [29]. Analogous reactions were observed in the case of pentose sugars [30]. As expected, aliphatic open-chain ketones are reduced at carbon, mercury, lead, or platinum electrodes mainly to secondary alcohols. The presence of a pro-chiral center in α-position leads to a mixture of erythro-threo isomers. Their ratio is discussed in terms of structure and solvent– electrolyte system [31].

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Organic Electrochemistry

Aliphatic cyclic ketones (e.g., cyclohexanones) being reduced offer a mixture of axial-equatorial isomers of the corresponding alcohols [32,33]. Reduction of ketosteroids in basic media gives rise, however, prevalently to equatorial alcohols [34]. 2. Aromatic Carbonyl Compounds In contrast to aliphatic carbonyls, the reduction of aromatic ones [35] (including alkyl-aryl ketones) leads to the pinacol as the main product, particularly in ethanolic or aqueous alkaline solutions on nickel, lead, or mercury cathodes. Already at the beginning of the twentieth century synthetic electrochemists were very active in this field. Pinacol can be, however, prepared also in acidic solution, but its yield and the ratio of (±) : meso stereoisomers differ from the alkaline reduction [36–38]. This finding indicates a different reduction–dimerization mechanism in acidic and in basic media, respectively. In acidic solution due to pre-protonation, only intermediate D is reduced, therefore dimerization involves two neutral radicals E. On the other hand, in basic environment the first reduction results in radical anion B (stabilized partly by conjugation with the present aromatic system), which is slowly (lack of protons) protonized to E. The latter species is then readily attacked by the excess of strongly nucleophilic radical anion B under the formation of pinacolate anion, which gets finally protonated. This mechanism prefers the formation of (±)-pinacol over the mesoform in basic media. The direct coupling of two radical anions should be in this case less probable due to charge repulsion [12]. The electroreduction of aromatic aldehydes in aprotic solvents was investigated in dimethylformamide (DMF) or acetonitrile (AN) using CV. The dimerization to pinacols was the main reaction process, and the reactivity sequence and properties of the radical species were evaluated [39]. Substituted benzaldehydes and acetophenones can be reduced in nonaqueous medium to pinacols also with phenol as a proton donor [40]. In very dry aprotic solvents, the only reaction pathway includes the radical anion B. Instead of its protonation, formation of ion pairs with cations of supporting electrolyte is possible. The eventual simultaneous interaction of two radical anions with one cation promotes sterically the (±)-pinacol. This effect was observed in acetonitrile or DMF with various cations like tetraethylammonium, sodium, europium(III) [41–43], and more recently also in reduction of alkyl-aryl ketones in an undivided cell with a magnesium or zinc sacrificial anode, where the generated Mg or Zn ions participate in ion pairing [44,45]. An indirect reduction appears sometimes to be very successful: 9-fluorenone in aprotic solvent is reduced in two reversible steps to anion radical and dianion, respectively (A–B–C). After preparative electrolysis at the first wave, only 10% of pinacol was isolated aside from 90% of the starting compound. After adding of titanocene (Cp2 Ti Cl2) as a complexing agent, the corresponding pinacol dominated (65–75%) among the products (A–B–E–H). At more negative potentials, however, the percentage of fluorenol increased and in the presence of a proton donor the alcohol was formed exclusively (A–B–C–F–G) [46]. A selective indirect reduction of aldehydes and ketones to alcohols can be achieved also in EtOH, with Al cathode and anode. The latter provides Al-ions, which form an intermediate Al(OEt)3 complex [47]. Pinacol, however, is not always the final product. At very negative potentials, the vic-diol can be further reduced under splitting of the C–C bond and the radical intermediates are protonated to alcohol (reduction H–G). Therefore, the reduction potential is the decisive factor for the proportion of pinacol versus alcohol. This phenomenon was observed, for example, in the case of benzaldehyde reduction in a membrane flow-cell using an acidic water–methanol mixture and a lead cathode [48] or indirect fluorenone reduction [46]. Several attempts have been done to the reduction of aromatic ketones in order to prepare optically active monoalcohols. For this purpose, an electrochemically generated chiral center (surface layer of assembled chiral compounds at the electrode–solution interface) or chiral conducting salt were used. However, the enantiomeric excess is still low and the results are rather irreproducible [8].

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Reduction of Aldehydes, Ketones, and Azomethines

Instead of dimerization leading to pinacols (type head-to-head), the reduced aromatic carbonyls in aprotic media can undergo also coupling of the type head-to-tail. This can occur due to the fact that the unpaired electron of the primary radical anion can be localized in various positions of the conjugated π-aromatic system, particularly in para-, alternatively in ortho-position relative to the carbonyl (according to the mesomeric scheme). As a result, carbonyl-to-ring coupling was observed, for example, in the case of 1-acetylnaphthalene in EtOH [49] or acetophenone [50]. Crossed pinacol coupling leading to otherwise difficultly available unsymmetrical pinacols (vic-diols) based on simultaneous electrolysis of aryl ketone and alkyl aldehyde accompanied by an intermolecular coupling (e.g., Reference 51) will be discussed in Section I.E.

C. REDUCTION OF CARBONYL COMPOUNDS INFLUENCED BY α-SUBSTITUTION In Sections I.B.1 and I.B.2, the reduction of compounds with an isolated carbonyl group could lead to alcohol, or to homo-coupling yielding a pinacol. In this section, a more complicated situation will be discussed, when the reduction of carbonyl is influenced by conjugation with an α,β-double bond or by substitution of the α-(β-) position, or when the neighbor double bond or some other grouping is activated by the carbonyl. It was already mentioned that the carbonyl group involves π-electrons that are able to interact with adjacent groupings under extension of the delocalized π-system due to mesomeric effects. If this interaction is strong enough, the molecule looses (partly or completely) the properties of individual original parts and exhibits a behavior characteristic for a new redox system. Simultaneously, carbonyls belong to the EWG. They provoke inductive shift of electron density in its neighborhood. The carbonyl carbon is thus partly positively charged; moreover, via the π-system of the double bond the β-position is also slightly positive (Scheme 31.4, upper four structures). Such an activated molecule can interact among others with the primary radical anion (father– son reaction), resulting in a cyclic hydroxycarbonyl product (Scheme 31.5). –0.51 O +0.48

+0.07

+0.39 H

–0.12

–0.50 O

–0.51 O

–0.43 O +0.07

–0.13

–0.19

–0.11

+0.10

+0.41

+0.18

+0.48

–0.51 O

–0.51 O +0.10 S-CH3

+0.51

0.00

–0.13

–0.33 O-CH3

–0.30 F

+0.46

+0.01

F +0.92

+0.12

SCHEME 31.4

–0.07 Cl

+0.50 –0.47 O

–0.51 O +0.49

–0.51 O

–0.27 NH2

+0.50

F

Selected atomic charges in various types of carbonyl compounds.

“Son”

O Ar

+e–

O O

Ar

“Father”

– +

+e–, +2H+

Ar

OH

O Ar

Ar “Father”

SCHEME 31.5

Example of the “father–son” reduction mechanism of α,β-unsaturated carbonyls.

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Organic Electrochemistry

In the case of α-substituted carbonyls, the significantly positive charge on the carbonyl carbon activates the adjacent bond and makes the reductive splitting of a substituent more favorable (Scheme 31.4, lower five structures), particularly when it is a good leaving group (selected atomic charges were calculated using [52]). As a result, the reduction of such activated carbonyl compounds does not necessarily lead to the reduction of the carbonyl group itself, but the β-position could be reduced first and an inversion of the expected sequence of reduction potentials may occur. In this case, the carbonyl acts as an activator, the C=O group is preserved in the first reduction step, and its reduction proceeds at more negative potentials. The composition of products and the individual mechanism then strongly depend on the particular structure of the parent compound and on the experimental conditions, particularly on the true working potential, solvent and cathode material. This situation, however, needs a more detailed comment: In the case of the electrochemical reduction of carbonyls influenced by α-substitution (this section), always a mixture of products is obtained; moreover, different dominating products or different proportions of isolated compounds were reported for the reduction of the same compound. How is it possible? The serious discussion of all possible mechanisms and the formulation of more general rules for the structure–reactivity relationship is very difficult because there are always two ultimate “target” questions, two aspects, two types of research teams, and thus two different experimental approaches and also two types of sources of experimental data: 1. From a fundamental point of view, the question is which part of the molecule is reduced first (= at less negative potentials), what are the primary intermediates (and their follow-up reactions), and what happens in the next reduction step at more negative potentials? For a detailed elucidation of the mechanism, voltammetric experiments on small electrodes with “analytical” (milimolar or sub-milimolar) concentrations are performed where direct heterogeneous reduction at the electrode may prevail (in connection with eventual adsorption of protonated species). The attention is focused particularly on the first reduction step and on the number of consumed electrons. The products of controlled-potential electrolyses are then determined by the working potential and thus various products at various potentials can be obtained and identified. 2. From an electrosynthetic point of view, the question is what are the final isolable products after complete reduction (without taking special care about the individual steps of the overall process). For this research, very often controlled-current electrolyses are performed on a preparative scale at large electrodes where the actual working potentials may reach very negative values. Several reduction steps thus can proceed simultaneously together with all accompanying chemical reactions. In addition to this, these reactions proceed at concentrations up to three orders of magnitude higher than in the previous case, hence the homogeneous processes—including coupling of generated radical species—play a significant role. It is necessary to stress that since the parent substances represent delocalized systems (particularly those bearing conjugated aromatic rings), the site of the electron attack may not be so important for the nature of the product, but rather the site of the subsequent protonation [4,53]. It should be also mentioned that if the γ-carbon is bearing a hydrogen, the latter is acidic and may influence the mechanism in aprotic media, particularly, when another EWG is attached to the γ-carbon. In the frame of this section, two examples will be demonstrated and discussed: α,β-unsaturated carbonyl compounds and molecules with a reducible substituent (often a good leaving group) in α-position toward the carbonyl. 1. α,β-Unsaturated Carbonyl Compounds Taking into account the generally known fact that the carbonyl function is reduced more easily (at less negative potentials) than an isolated double bond, one can expect that in molecules bearing

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Reduction of Aldehydes, Ketones, and Azomethines

O

+e–

O–

O–

O – Tail

Head O

SCHEME 31.6

+e–, +H+

OH

O

Mesomeric structures of the first reduction intermediates.

simultaneously carbonyl and a double bond, the C=O group will be always reduced first. However, in α,β-unsaturated carbonyls, the conjugated system –CO–C=C– changes the situation: the same molecule can be considered either as an aldehyde (ketone) with conjugated olefin or as an activated double bond. The presence of an aromatic ring attached either at the α-carbon adjacent to the carbonyl or at the β-carbon of the α,β-double bond causes further extension of the delocalized π-system. Due to this, the reduction center (LUMO) is not located exclusively at the carbonyl, but involves a larger part of the molecule and the primary radical anion can exist in several mesomeric structures (Scheme 31.6). Therefore, the electrochemical reduction of α,β-unsaturated aromatic carbonyls may result not only in the reduction of the carbonyl to alcohol or in the reductive coupling to pinacol, but also to the exclusive saturation of the α,β-double bond. It is necessary to mention that many coupling products form various stereoisomers. Their detailed discussion is, however, beyond the scope of this chapter. a. Aliphatic and Alicyclic Carbonyls Taking into account Scheme 31.6, after the first one-electron reduction the unpaired electron is localized either prevalently at the carbonyl carbon (head position) or at the β-carbon (tail position). These two radical species are able to undergo radical coupling. Theoretically three combinations can occur: a tail-to-tail coupling, when δ-diketones are formed, head-to-head hydrodimerization yielding a pinacol, and head-to-tail coupling resulting in γ-hydroxyketone. Very often a mixture of all products is obtained. The predominant mechanism is structurally controlled. A freely accessible β-position or sterical hindrance close to the carbonyl may give preference to tail-to-tail coupling; on the other hand, a hindered β-position can favor the head-to-head link. The classic examples are reduction of mesityl oxide [54] or ionone [55]. According to the reaction conditions, the primary coupling products may undergo follow-up reactions giving rise to cyclic final products (Scheme 31.7). To increase the proportion of the pinacol-type coupled products, the presence of Sn(II) or Cr(III) ions in the solution has a positive effect. Due to their coordination with primary anion radicals through the oxygen atom, the radical character of the intermediates is accentuated and the pinacolization is preferred [42]. Similar reaction pathways were observed in the case of cyclohexenone [56,57]. For the nonsubstituted parent compound (R1, R2, R3 = H), exclusively tail-to-tail products were formed. For R1, R3 = Me, R2 = H or R2, R3 = Me, R1 = H, preferentially head-to-head and head-to-tail derivatives were produced. Benzalacetones should be regarded as intermediate case between aromatic and aliphatic carbonyls because the aromatic ring is not attached directly to the carbonyl. The final products after galvanostatic reduction in acetonitrile are cyclic hydrodimers. The mechanism of their formation follows the discussed principles: the first electron attacks the double bond under formation of a radical with unpaired electron at the β-carbon. The tail-to-tail coupling results in a linear hydrodimer that undergoes cyclization by attachment of the partly positive carbonyl carbon to the negatively charged α-carbon (Scheme 31.8). Two stable diastereoisomers were isolated [58].

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Organic Electrochemistry Head-to-tail O

+2e– +2H+

Tail-to-tail

O

+

Hcetate buffer Hg-cathode

2 Mesityl oxide

HO O

O OH CH3 O

CH3 H

O O

O

OH OH

R1

+2e– +2H+

2

AN/acetate buffer or DMF/TBAI

Ionone

R3

R3

Head-to-head

R2

Cyclohexenone

SCHEME 31.7 Examples of “tail-to-tail,” “head-to-tail,” and “head-to-head” reductive coupling of α,βunsaturated carbonyls. O

O

OH +H+ +e– – O – C3 C1 C2

O

O

Dim. +H+

+H+

OH

O

SCHEME 31.8 Electroreduction pathway of benzalacetones.

The aforementioned electroreductive hydrocoupling of benzalacetone (and also of methyl cinnamate, methyl crotonate, and coumarin) was studied by DFT calculations (B3LYP/6–311++G**). [15]. The Mulliken spin densities and atomic polar tensor (APT)–derived charges for C1, C2, and C3 of the primary radical anion (see Scheme 31.8) are shown in Table 31.1. The calculated data for transition states correspond well to the experimental results and confirm the suggested mechanism. TAbLE 31.1 Spin Densities and Atomic Charge for Carbonyl Carbon (C1), α-Carbon (C2), and β-Carbon (C3) of the Primary Radical Anion of benzalacetone Carbon C1 C2 C3

Spin Densities

Atomic Charge

0.12 0.08 0.27

+1.18 −0.66 −0.11

Source: Kise, N., J. Org. Chem., 71, 9203, 2006.

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Reduction of Aldehydes, Ketones, and Azomethines Aryl

O

O

O

Aryl

O

O

Aryl

Aryl

O Aryl

Aryl +

trans-

SCHEME 31.9

cis-

Electrochemical reduction of bis(α,β-unsaturated arylketones). O X

O +

Alkyl

Electroredn. DMF/AN 1:1, Ni-cath.

Alkyl

SCHEME 31.10 Electrochemical reduction of methyl vinyl ketone in the presence of an aromatic halide; X = halogen atom.

A special example of this type of reaction is the electrochemical reduction of bis(α,β-unsaturated arylketones) [59] discussed later. The coupling (intramolecular) does not proceed like in Schemes 31.7 or 31.8, but symmetrically under the formation of a diketone (Scheme 31.9). The reduction of methyl vinyl ketone in the presence of an aromatic halide is a very selective electrosynthetic method for arylation of the carbonyl compound. It was used as a key step in preparation of a precursor for the synthesis of medium-ring benzolactones [60]. The unpaired electron of the primary radical anion is localized at the activated vinyl double bond (see Scheme 31.6), but instead of the expected homocoupling tail-to-tail, the vinylic radical reacts with the halogen. Hence, the reductive coupling of the vinyl to the aromatic system occurs under splitting of the halide ion, whereas the carbonyl function remains preserved (Scheme 31.10). b. Aromatic Carbonyls Whereas in aliphatic and alicyclic α,β-unsaturated carbonyls the reductive coupling prevails, in aromatic α,β-unsaturated carbonyls the saturation of the double bond is more typical. Aromatic carbonyls are primarily those bearing the aromatic ring directly at the carbonyl carbon (derivatives of benzaldehyde or acetophenone). Cathodic reduction of phenyl vinyl ketone [61,62] and chalcone [63] in buffered aqueous solutions is the typical example of inversion of the expected sequence of reduction processes: at first the two-electron two-proton reduction of the double bond occurs yielding the corresponding alkyl aryl ketone, which is at more negative potentials reduced to alcohol. Similarly, aldehydes like cinnamaldehyde [64] due to the presence of an aromate are reduced to saturated aldehydes, whereas the reduction of alkenyl aldehydes (like crotonaldehydic ring) results in the unsaturated alcohol [65]. An electrolysis of cinnamaldehyde in buffered aqueous media gave a mixture of 3-phenylpropionaldehyde and cinnamyl alcohol [66]. Even in this case, the first reduction step involved saturation of the double bond, the cinnamyl alcohol is probably the product of reduction of pinacol at very negative potentials (see Figure 31.1, reaction H−G). More recently, a series of chalcones (phenyl styryl ketones) substituted at both phenyl rings was investigated in nonaqueous DMF. Besides the evaluation of multiple substituent effects (see Section I.A), it was concluded that two successive one-electron steps for the reduction of the chalcones yielded always saturation of the double bond [14]. Chalcones and their mesityl- and anisyl- analogs were reduced also in aqueous butanol [13]. In all cases, reduction of the double bond is dominating. In the mixture of products, the linear δ-diketones formed in the tail-to-tail coupling were isolated together with their cyclic form (see Scheme 31.8).

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Organic Electrochemistry

2. α-Substituted Carbonyls O Y Y = –NR2, –OH, –SR, –F, –Cl, or Br

Like in Section I.C.1, it should be kept in mind that the isolated carbonyl group is electrochemically reduced less negatively (more easily) than the isolated C−Y bond of organic species bearing groups Y like –NR 2 , –OH, –SR, or halogen. The latter substituents represent leaving groups able to undergo a reductive elimination. The situation is changed when the group Y is attached to the α-carbon adjacent to the carbonyl group. The electron-withdrawing ability of C=O causes activation of the C–Y bond, which is reduced first contrary to expectations. The carbonyl remains unchanged and is reduced at more negative potentials. This effect is described for glycolaldehyde [67], α-aminoketones [68], phenacyl sulfonium ions [69,70], or α-halogen-substituted carbonyl compounds [71]. In the case of electrochemical reduction of ω,ω,ω-trifluoroacetophenone, in acidic media below pH 5, the protonated carbonyl is reduced to pinacol and alcohol. In neutral and basic pH range, however, all C–F bonds are reductively split off prior to the reduction of carbonyl [72]. Besides the pre-protonation, the hydration of the carbonyl group at higher pH values (preventing its reduction) plays also an important role [73,74] The electroreductive defluorination of trifluoromethyl ketones was utilized for electrosynthesis of various fluorinated derivatives in acetonitrile at a lead plate cathode in the presence of chlorotrialkyl silane. Among others, 2,2-difluoroenol silyl ethers were obtained [75,76]. A classical case of inversion of the expected sequence of reduction processes caused by mutual electronic interaction of a carbonyl with an unsaturated substituent is represented by benzaldehydes substituted in the para- (or ortho-) position by a –CN group. It is generally known that the electrochemical reduction of a nitrile is more difficult than carbonyl reduction. However, even here the carbonyl forms a π-electron delocalized system with the nitril via the phenyl ring. As a result, in acidic media (pH < 3.5) the protonated nitrile group is reduced first by four electrons to an amine, whereas the carbonyl is reduced at more negative potentials [77]. Electrochemical reduction of aliphatic carbonyl compounds activated by a nitrile function is a bit different and more complicated. In the reduction of 2-cyanocycloalkanones (Scheme 31.11), the reaction pathway strongly depends on the other α-substituent R and on the nature of the cation of supporting electrolyte. When R = H, the carbonyl itself is reduced and the corresponding cisβ-hydroxynitrile is formed. When, however, R = alkyl, and the supporting cation is large (e.g., Et4N+), cleavage of the C−CN bond occurs. The role of the cation consists in stabilization of the primary radical anion by ion pairing [78].

CN

CN +2e–, +2H+ OH

O R

–

R CN

+e–

R CN

+e–, +H+ –CN–

O

SCHEME 31.11

O

Electrochemical reduction of 2-cyanocycloalkanones.

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O

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Reduction of Aldehydes, Ketones, and Azomethines

D. REDUCTION OF CARBONYLS WITH A REMOTE UNSATURATION–INTRAMOLECULAR COUPLING 1. Remote Double bond or Aromatics Reduction of a carbonyl compound with a remote double bond or an aromatic ring can be accompanied by intramolecular coupling of the primary radical anion with the double bond under closure of a new ring. The C–C bond formation occurs on the more substituted carbon of the double bond, closer to the carbonyl group (Scheme 31.12). The cyclization leading to a five- or six-membered ring is the most effective, although the formation of four- as well as seven-membered rings is also possible. The reaction proceeds usually in mixed organic solvents—mostly in MeOH/dioxan or 2-propanol/DMF. The cyclization is regio- and stereoselective [79–82]. Sometimes, however, the direct reduction of the carbonyl for example, in DMF yields simply secondary alcohol without cyclization (Scheme  31.13). The proportion of cyclic versus acyclic forms depends strongly on the cathode material, solvent, and supporting electrolyte [83]. Upon addition of N,N′-dimethylpyrrolidinium tetrafluoroborate [84] or N,N′-dimethylquininium tetrafluoroborate as a mediator and cyclizing agent controlling stereoselectivity, cis-cyclized products are obtained in good yield [85]. The mechanism was often discussed [86] and recently reconfirmed by ab initio as well as DFT calculations [87]. After the first electron transfer, the spin density of the primary radical anion is localized at the C2 atom, but during the progress of intramolecular coupling the terminal double bond is approaching the carbonyl moiety (see Scheme 31.12) and the spin density is shared between C2 and C6+C7. Based on optimization of various transition states, the comparisons of the relative activation energies indicate that the regio- and stereoselectivities in the cyclizations of the ketyl radicals are determined by kinetic control. In the case of aromatic nonconjugated ketones, the intramolecular cyclization may occur when the chain between the carbonyl and phenyl groups involves three (possibly also two or four) carbon atoms, resulting in five- to seven-membered ring. It was proven that the presence of methyl groups in positions C3, C4, or C5 (see Scheme 31.13) does not prevent the cyclization [83]. However, the second substitution on the phenyl ring is important for this process: electron-donating substituents (Me, OMe,…) in meta- position prevent the cyclization, whereas an EWG like CN or COOAlk in meta-position (and/or Me, OMe in para-position) does not inhibit it. This finding points to the interpretation that the reactive species attacking the aromatic ring has an anionic character. OH

O O C4

C6

Electroreduction

C2

CH3

C1 C7

C5

C3

C1

CH3 C7

OH

O

O

Electroreduction

CH3 H

SCHEME 31.12 Ring closure during electrochemical reduction of a carbonyl compound with a remote double bond. OH parameta-

O C4 ortho- C5

O

Electroreduction

H

CH3

HO

+ C3

SCHEME 31.13 Ring closure and alternative secondary alcohol formation during electrochemical reduction of a carbonyl compound with remote aromatic rings.

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Organic Electrochemistry

+e– O

N

H OH

+e–, +2H+ O–

N

N

trans-cyclization

I

O — C–OMe

O

O–

O

H

+e–, +2H+ trans-cyclization

N

O

+e– N

O

N

O — C–OMe

O

H

+e–, +2H+ cis-cyclization

II

OH

N

OH

SCHEME 31.14 Stereospecificity in the reduction of 1-indolealkanones (I) and 3-methoxy-carbonyl-1indolealkanones (II).

Another example of how substitution on the aromatic ring can influence the resulting stereospecificity is reported for 1-indolealkanones (I) and 3-methoxycarbonyl-1-indolealkanones (II) in isopropanol [88]. The observed trans- and cis-cyclized products are formed according to two different mechanisms: whereas the unpaired electron in the anion radical of I is located as usually at the carbonyl carbon, in the case of II the strong electron-withdrawing effect of the ester group in position 3 changes the reduction center to the indole ring and thus the methoxycarbonyl group is reduced more easily than the carbonyl. The attacking radical is located at a different place causing the change of mechanism and thus stereospecificity. This interpretation was supported by DFT calculations where the highest spin density was found at position 2 of the heterocycle (Scheme 31.14). When the linear chain of nonconjugated enones contains a nitrogen or sulfur heteroatom, this is the way how to prepare various saturated heterocyclic alcohols. Nevertheless, the nitrogencontaining derivative exhibits some decrease in stereoselectivity in electroreductive intramolecular cyclization in comparison with high stereoselectivity for the analogous enones possessing a sulfurcontaining chain or an all-carbon chain [89]. 2. Remote Unsaturated groups An example of intramolecular coupling of a carbonyl with a remote reactive group is the cathodic reduction of aromatic ketones with an imino-ether group in the ortho-position. In acetonitrile, the carbonyl is reduced first and undergoes a cyclization under formation of a benzoxazine. In basic solution, a quinoline is formed (Scheme 31.15). The presented electrosynthetic reaction can be an alternative way to benzoxazines and quinolines [90]. R1

H O

O

AN

R2

N R1

R2

Electroreduction

N

R1 O–R

R2

Base N

SCHEME 31.15 group.

O–R

Reductive intramolecular coupling of an aromatic carbonyl with a remote unsaturated

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1217

Reduction of Aldehydes, Ketones, and Azomethines OH

O +e, Sn cath. CN 2-PrOH O

+e, Sn cath. 2-PrOH

No reaction

H2O

CN 2-PrOH

SCHEME 31.16 γ- or δ-position.

O

OH O

OH NH +e, Sn cath.

+

Electrochemical reduction of a saturated carbonyl compound bearing a nitrile group in the

Electroreductive cycloaddition of symmetric bis(α,β-unsaturated arylketones) leads among others to cyclobutane derivatives (Scheme 31.9). Diastereoselectivity of this reaction is strongly controlled by the nature of the electrolyte cation. The formation of cis-cyclobutanes is favored by Mg or Ba ions, probably due to a chelation effect [59]. Cathodic reduction of a carbonyl molecule bearing a nitrile group in the γ- or δ-position is accompanied by the intramolecular cyclization resulting in an α-hydroxy ketone. It was assumed that the attacking reduced carbonyl has a radicalic nature rather than anionic since the reaction proceeds in protic media (2-propanol). The reductive addition of the primary carbonyl radical gives rise to a cyclic iminoalcohol, which is hydrolyzed (during the workup) to the α-hydroxyketon and its dehydroxylated derivative as the major products [91] (Scheme 31.16). A more recent theoretical study of reductive intramolecular coupling of carbonyl compounds offers a more detailed explanation based on sharing of the unpaired electron upon sterical approach of the unsaturated group to the primary carbonyl radical (anion) between these two—see Reference 87. Also, the electroreductive coupling of a carbonyl group with an intramolecular O–Me oxime moiety gave the corresponding cyclized product stereoselectively [92].

E.

REDUCTION OF CARBONYLS IN THE PRESENCE OF ANOTHER REACTANT–INTERMOLECULAR COUPLING

In this section, the reported reactions are often analogous to those of the previous part concerning intramolecular processes; the possible differences are caused by lower sterical requirements. When a mixture of aryl ketone and alkyl aldehyde or ketone is electrolyzed in the presence of chlorotrimethylsilane (CTMS), a crossed pinacol coupling occurs, leading to otherwise difficultly available unsymmetrical pinacols (vic-diols) [51]. This type of rather general reactions starts with the electroreduction of the aromatic carbonyl to the corresponding radical anion. Its immediate reaction with CTMS present in the solution causes indirect reductive splitting of chloride anion under formation of a silyl derivative with “umpolung” of the carbon–oxygen bond (see Scheme 31.17). The nucleophilic carbanion adds to the electrophilic center of the still nonreduced carbonyl partner (alkyl ketone, acyl or ester) and an unsymmetric pinacol is obtained. When the second carbonyl partner contains a good leaving group, for example, imidazol (Im), in the case of intermolecular electroreductive coupling of aromatic ketones with 1-acylimidazoles, an α-hydroxy ketone is formed [93]. The aromatic carbonyls are first reduced in the presence of CTMS and triethylamine (TEA), after the coupling with the acylimidazole, the desilylation using TBAF in THF follows (Scheme 31.17). This reaction is effective also for the intramolecular reductive coupling of γ-, δ-, ε-, and ζ-keto acylimidazoles, resulting in four- to seven-membered rings [93]. Reductive allylation of carbonyls represents an important way to homoallylic alcohols. Electrolysis of a mixture of an aldehyde or ketone with an allylic halide in hexamethylphosphor-amide (HMPA)

© 2016 by Taylor & Francis Group, LLC

1218

Organic Electrochemistry O +

+2e– +CTMS CH3

R1

R2

CH3 R1

– CH3

–

–Cl

R2

TBAF/THF

O

OH OH

TMS-O O

+ H3C

lm

–Im–

CH3 O–

CH3 O

lm TMS-O

CH3

TMS-O

CH3 O

TBAF THF

OH

CH3

CH3

SCHEME 31.17 Electrochemical reduction of an aryl ketone in the presence of a saturated carbonyl compound.

results in the desired β,γ-unsaturated alcohol. The mechanism of this reaction and the regioselectivity depends, however, on the more readily reduced reaction partner. In the case of aryl aldehydes or alkyl aryl ketones being reduced first, the primary radical anion attacks the halide in an SN2 fashion (Scheme 31.18). In the case of aliphatic carbonyls (acetone, propanal), the allylic halide is primarily reduced by two electrons, giving rise to an allylic carbanion that adds at the electrophilic carbonyl carbon (Scheme 31.19). The influence of the electrode material on the yield is crucial [94]. An analogous electrochemical allylation with high efficiency has been performed in mixed aqueous acidic and basic media on a Zn cathode with Zn ions as catalysts [95]. Cathodic intermolecular coupling of aliphatic ketones with 1-olefins (or dienes or trienes) leading to the corresponding tertiary alcohols has been satisfactorily accomplished with high regioselectivity by using a carbon fiber cathode in DMF [96] (Scheme 31.20). Coupling of ketones with dienol ethers under the same conditions (DMF, carbon fiber), however, has only a very low yield. For the successful coupling, the addition of CTMS and TEA into the O Ar

O–

+e– Ar

R2 O–

R2

R4 +

Ar

SCHEME 31.18

R2

R3

Cl

–Cl–

R4

OH

SN2 : +e–, +H+

R3 Ar

R2

Reductive allylation of aromatic carbonyls resulting in homoallylic alcohols. R4

+2e–

R4

–Cl–

–

R3

Cl

R3

R4

O

–

Addition; +H+

+ R3

R4 R3 OH

R1

R2 R1 R2

SCHEME 31.19

Electroreduction of the mixture of a saturated carbonyl compound with an allylic halide.

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1219

Reduction of Aldehydes, Ketones, and Azomethines O

O–

+e– R2

R1

R1

OH

O–

SCHEME 31.20

+e–; +2H+ R2

R1

R2

R3

+

R1

R3

R2

Cathodic intermolecular coupling of aliphatic ketones with 1-olefins.

O–Me

O-Me

O el. redn.

+

OH

O

OH

H3O+

CTMS/TEA O Me–O

SCHEME 31.21

OH el. redn. CTMS/TEA

+

H3O+

OH O

Reductive coupling of ketones with dienol ethers.

O R1 O

+

H

CH3NO2

O

Reductive electrocatalysis MeOH

Reductive electrocatalysis + CH3NO2

SCHEME 31.22

Me-O

OH R1

NO2 H O HO NO2

MeOH

Electroreduction of the mixture of carbonyl compounds and nitromethane.

reaction system was necessary. The following acidic hydrolysis leads to the corresponding remote hydroxy keton or aldehyde [96] (Scheme 31.21). The electroreduction of the mixture of carbonyl compounds and nitromethane in methanol yields the corresponding β-nitro alcohols according to the electrochemically induced Henry reaction (Scheme 31.22). This process is an example of an electrocatalytic system. The electrochemical reduction of the reaction mixture (under galvanostatic conditions) yields an electrochemically generated base catalyzing the desired reaction. The detailed mechanism depends on the composition of the solution [97]. Electroreduction of a mixture of various ketones with aliphatic O-methyl oximes gives rise a β-methoxyamino alcohols that can be easily (chemically) reduced to β-amino alcohols (Scheme 31.23). A similar mechanism operates in the case of the electroreductive intermolecular coupling of ketones with N,N-dimethylhydrazones or nitrones [92]. Analogously to previous reactions, when aldehydes are reduced together with phthalimides (in the presence of CTMS and TEA) intermolecularly coupled products (3-hydroxy-3-(1-hydroxyalkyl)isoindolin-1-ones) are formed as an important precursor for further stereospecific syntheses [98] (Scheme 31.24). When aromatic aldehydes are electrolyzed in the presence of aryldiindenylmethanes, the monoarylidene derivatives are formed [99]. Intermolecular reductive coupling of carbonyl compounds with nitriles results in an α-hydroxy carbonyl derivative that is equivalent to the product of acylation [93]. The suggested mechanism is based on the fact that the unactivated carbonyl is reduced more easily than a nitrile. As the

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1220

Organic Electrochemistry O

O–

+e–

R1

R2

R1

O–

R2

N–O–CH3 +

R1

R2

R3

NH–O–CH3

HO

+e–, +2H+

R1

R4

R2

HO Redn.

R4

NH2

R1

R3

R4 R2

R3

SCHEME 31.23 Electrochemical reduction of various ketones with aliphatic O-methyl oximes leading to β-amino alcohols.

O O N X

+ R

R

1) +2e– CTMS/TEA

HO

OH N

2) Bu4NF

H

O

X

O

SCHEME 31.24 Electrochemical reduction of aldehydes with phthalimides.

O R1

+ R3(R4)CH–CN

OH

Electroredn. Sn-cathode

R2

OH R4

O

R1

+ R1 R2

CH(R4)R3

CN R2 R3

SCHEME 31.25 Formation of α-hydroxy carbonyl derivatives in the intermolecular reductive coupling of carbonyl compounds with nitriles. COOH O

CH3

HO +

O

C

O

CH3

Electroredn. DMF

R

R

SCHEME 31.26 Preparation of aromatic α-hydroxy acids using electrochemical reduction of aromatic carbonyl compounds in the presence of carbon dioxide.

by-product, a β-hydroxy nitrile is formed. Its exclusive production occurs when only aprotic nitriles serve as a solvent (acetonitrile or other alkylnitriles). When, however, a certain proportion of an alcohol (i-PrOH, EtOH) is added, due to the protonation of the primary radical anion, a radical addition takes place, leading (after hydrolysis of the imine intermediate) to the desired α-hydroxy carbonyl [91] (Scheme 31.25). Electrochemical reduction of aromatic carbonyl compounds (in DMF) in the presence of carbon dioxide results in a carboxylated product. This process was used for synthesis of various aromatic α-hydroxy acids [100–104] (Scheme 31.26). Potential controlled reduction of a mixture of aldehydes or ketones with a 10 to 15-fold excess of primary amines yields secondary amines. The reaction proceeds at a Hg cathode in the aqueous solution of the amine, half-neutralized with HCl, hence in buffered media of pH 10–11. The detailed sequence of electron and proton transfers was not reported, most probably the pathway involves the corresponding Schiff base as intermediate, which is reducible, giving rise to the product (Scheme 31.27). In the case of cyclic ketones, the reaction is diastereoselective [105].

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1221

Reduction of Aldehydes, Ketones, and Azomethines R1

Controlled potential reduction O

+

H2N–R3

R2

SCHEME 31.27

Hg, pH 10-11

R1

H N

R2

R3

Cathodic reduction of aldehydes or ketones with an excess of primary amines.

F. REDUCTION OF DICARBONYL COMPOUNDS The dicarbonyl compounds can be divided into several categories: aromatic ones, where in o- and p-dicarbonyls, the C=O groups are directly involved in the aromatic π-system (in contrast to m-dicarbonyls, where the electronic communication between the two carbonyls is negligible), α- and β-dicarbonyls on an aliphatic chain, where the two close C=O groups influence each other, and “remote” dicarbonyls where the C=O groups are wide apart with lacking delocalization between them—hence, practically independent (in fact, aromatic m-dicarbonyls belong to this type). The main feature in the first two categories is that one of the C=O groups plays the role of a strong EWG toward the other one. There are two main consequences: (1) the reduction of the first carbonyl is shifted to less negative potentials due to the effect of the second one; (2) owing to the fact that hydration is enhanced by the influence of an EWG, in aqueous media those dicarbonyls are strongly hydrated under the formation of a geminal diol, which is electrochemically inactive. One should also keep in mind that also keto-enol tautomery can play its role. Therefore, the reduction of dicarbonyl compounds is very often complicated by simultaneous acidobasic, hydration/dehydration, and tautomeric equilibria. 1. 1,2-, 1,3-, and 1,4-Dicarbonyls on an Aromatic Ring The most simple dialdehydes where both carbonyl groups are attached to a benzene ring (or more generally to an aromatic system, e.g., naphthalene, etc.) can exist as three isomers: 1,2-benzenedialdehyde (phthalaldehyde, orthophthalaldehyde, OPA), 1,3-benzenedialdehyde (isophthalaldehyde, IPA), and 1,4-benzenedialdehyde (terephthalaldehyde, TPA) [4]. In 1,2- and 1,4-isomers, the carbonyls are in conjugation, their mutual influence is substantial, and they behave as benzaldehydes with a strong EWG substituent. In contradistinction to this, the 1,3-isomer has an interrupted π-system between the two C=O groups. Taking into account that the hydration reaction is promoted by the π-interaction with EWG substituents, OPA is hydrated very strongly—approx. 85%, TPA is also considerably hydrated (about 15%) whereas IPA is hydrated only negligibly (up to 3%, like benzaldehyde). As a result, the reduction of IPA proceeds analogously to benzaldehyde (with double current) [106]. The reduction of TPA (and other related p-dicarbonylaromatics) in acidic solutions is analogous to the noteworthy mechanism of p-diacetylbenzene reduction, observed already in the 1960s, which exhibits in this media a rare two-electron reversible process [107]: the diprotonated form accepts simultaneously and reversibly two electrons giving rise to a tautomeric system where the quinonemethide with limited stability is the most important intermediate that is at more negative potentials reduced to diol (the biradical species in Scheme 31.28 was not proven yet). The same mechanism was recently found and confirmed by quantum chemical calculations in nonsymmetric analogs of such p-phenylene dicarbonyls where a carbene or imine moiety is located in para-position with respect to the carbonyl [24]. The strong hydration of OPA (particularly at pH 2–6) is further complicated by steric reasons (ortho-effect), resulting in two successive hydration equilibria: besides the simple addition of water to the carbonyl yielding a geminal diol, a follow-up cyclization occurs giving rise to electrochemically inactive isobenzofurane-1,3-diol (Scheme 31.29). The hydration/dehydration kinetics is strongly dependent on pH, hence at pH 2–6 the OPA is effectively nonreducible [108]. The complete electroreduction of these three benzene dialdehyde isomers in buffered solutions yields the corresponding diols, the detailed mechanism and sequence of electron and proton transfers depend on pH [108–110].

© 2016 by Taylor & Francis Group, LLC

1222

Organic Electrochemistry O

R

R

OH

R = H, CH3 ,C6H5

R

O

R

OH

+2H+ OH+

R

R

R

OH

+2H+ +2e–

+2e–

OH+

R

SCHEME 31.28

R

C R H OH

C R H OH

OH

Electrochemical reduction of p-diacetylbenzene. OH

O +H2O

OH

OH

O

SCHEME 31.29

R H OH C

O

O OH

O

Hydration equilibria of phthalaldehyde.

2. 1,2-Dicarbonyls A number of papers devoted to the electrochemical reduction of aliphatic (e.g., 2,3-pentanedione), alicyclic (1,2-cyclohexanedione), and aromatic (1-phenylpropane-1,2-dione or 1,2-diphenylethane1,2-dione = benzil) α-dicarbonyl compounds in buffered aqueous solutions were reported [111–116]. Besides the dicarbonyl form, the parent compounds may exist in two other forms: as an enolic species and in hydrated form due to the close vicinity of the other carbonyl (Scheme 31.30). Besides, protonation (or diprotonation in acidic media [113]) may precede the electron transfers, whereas in neutral and basic solution the first electron transfer is (quasi)reversible [115]. OH+ O +2H+ OH

OH+ OH+

O

OH

+H+ O

+2e–

+2e– +H+

Slow

O

–H2O O +H2O

HO OH

O +e–

+e– +2H+ O–

OH

OH +2e– + OH +2H

O OH

SCHEME 31.30 Tautomeric, acidobasic, hydration, and redox transformations of 1,2-dicarbonyl compounds.

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1223

Reduction of Aldehydes, Ketones, and Azomethines

The totally two-electron, two-proton reduction yields at first the enediol, which is oxidizable at less negative potentials (the proof of its presence can be carried out by CV). The enediol is then transformed into the α-hydroxy carbonyl (ketol), which is the first stable product for the majority of α-diketones, particularly aliphatic and alicyclic (Scheme 31.30). At more negative potentials, aromatic ketols can be further reduced to diols [113]. A typical example is benzil: in acidic-neutral buffers up to pH 8, the reaction sequence involves pre-protonation (H+, e –, e –, H+); in more acidic solutions antecedent diprotonation occurs prior the two-electron reduction (Scheme 31.30). The follow-up isomerization is surface-assisted, and the trans-cis interconversion is generally slower than the cis-trans. The resulting enediol then slowly tautomerizes into benzoin, which is again reducible by two electrons and two protons at more negative potentials giving rise to a diol. A recent study involves a detailed description of kinetics and equilibria of the mentioned process in the H0/pH region −5 to 11 [117]. This mechanism is similar for both alkyl and aryl dicarbonyls as well as for diketones and ketoaldehydes (like methylglyoxal). The minor differences are (a) in hydration, which is more typical for alkyl dicarbonyls; (b) in protonation, which proceeds first at the aldehydic carbonyl (in ketoaldehydes) or at the aromatic one (in alkyl-aryl dicarbonyls [113]); (c) in tautomeric forms, where the enol form is predominant in cyclic diones; and (d) in regioselectivity, which is particularly in asymmetrical aliphatic diketones poor. Although the hydroxyketone is the main electroreduction product of alkyl diketones in neutral media, the aliphatic and aldehydic species may also undergo (as a side reaction) a reductive dimerization to diastereomeric keto-pinacols [116]. The alternative way from diketones to stereochemically pure diols is also a microbial reduction combined with electrochemistry [116]. The correlation of reduction potentials in a series of dicarbonyls and their analogs with structure and bioactivity has been also discussed [118]. In aprotic solvents, the first step of electrochemical reduction of α-diketones proceeds reversibly. In the presence of metal ions (group IA and IIA), ion pairing with the primary radical anion was observed [119,120]. 3. 1,3-Dicarbonyls and Remote Diones 1,3-Dicarbonyls exist in solution in equilibrium with their enol form, which is stabilized by hydrogen bonding as a six-membered ring (Scheme 31.31). In the case of aromatic species (e.g., dibenzoylmethane), the system is further conjugated with the two aromatic ring systems [121]. The hydrogen atoms are labile and can be replaced by metal ions giving rise to a chelate ring [122]. Many derivatives with various combinations of electron-withdrawing and electron-donating substituents were electrochemically investigated and the results were successfully correlated with DFT calculations [123]. The electroreduction mechanism was proposed already several decades ago: in aprotic solvents, 1,3-diphenyl-1,3-propanedione (dibenzoylmethane) is reversibly reduced by one electron to an enolate radical anion, which is protonated by proton transfer from the parent molecule. The formed radical then dimerizes to pinacol. As a result, in total a “half-electron” process occurs (Scheme 31.32) [124]: In the case of dicarbonyls separated by a longer aliphatic chain (three carbon atoms and more) where no mutual interaction between them takes place, the cyclization like in the section on carbonyls with remote unsaturated bond occurs. As an example, the reduction of 1,3-dibenzoylpropanes— III (1,5-dicarbonyls) gives rise to a 1,2-cyclic diol—IV (Scheme 31.33). O

SCHEME 31.31

O

Tautomeric equilibrium of 1,3-Dicarbonyls.

© 2016 by Taylor & Francis Group, LLC

O

H

O

1224

Organic Electrochemistry H O

O H

+2e–

O

O

4 Ph

O

+ 2 Ph

Ph

O–

Ph

Ph

Ph

Ph O

Ph

O H

SCHEME 31.32

The formally “half-electron” process of 1,3-diphenyl-1,3-propanedione reduction. O

O

OH OH +2e–, +2H+

R

R III

SCHEME 31.33

a: R = H b: R = phenyl

R

R IV H2

Intramolecular cyclization during electroreduction of 1,3-dibenzoylpropanes.

As for the mechanism, two pathways were suggested: either the cyclization occurs on the level of dianion of the parent compound III (Equations 31.1 through 31.4) [125,126] or on the level of radical anion (Equations 31.1 and 31.4 through 31.6) [127]. III + e− ⇆ III− •

(31.1)

(a) 2 III− • → III2− + III III2− → IV2− (cyclization) IV2− + 2 H+ → IV H2

(31.2) (31.3) (31.4)

(b) III− • → IV− • (cyclization)

(31.5)

IV− • + III− • → IV2− + III

(31.6)

IV2− + 2 H+ → IV H2

(31.4)

More recently, the electroreduction of aryl 1,4-, 1,5-, or 1,6- (γ-, δ- or ε-) diketones in the presence of CTMS was studied from the stereospecific electrosynthetic point of view and the formation of unsymmetrical cyclic pinacols (vic-diols) via intramolecular head-to-head coupling with transstereoselectivity was confirmed (Scheme 31.34). It is notable that in contrast to electroreduction, the chemical reduction using TiCl4 –Zn yields prevalently cis-isomers [51,128]. An analogous reaction with γ-, δ-, ε-, and ζ-keto acylimidazoles [93], however, gives rise to cyclic α-hydroxy ketones (instead of cyclic pinacols) due to the leaving imidazole anion (Scheme 31.35). 4. β-Keto Esters In the electrochemical reduction of β-keto esters, the carbonyl is reduced first, therefore this topic was inserted in the present chapter. This process, accompanied by the Tafel rearrangement, is already a “classic” reaction [8]. The accepted reaction mechanism is the following [129]: under acidic

© 2016 by Taylor & Francis Group, LLC

1225

Reduction of Aldehydes, Ketones, and Azomethines O

OH OH Electroreduction

Ph

Ph

X

X

O

X = (–CH2–)1–3 HO

O

O

HO

Y

Y

Electroreduction

CO-OEt

CO-OEt Y = (–CH2–)1–2

Electrochemical reduction of aryl 1,4-, 1,5-, or 1,6- (γ-, δ- or ε-) diketones.

SCHEME 31.34

O

Electrochemical reduction of γ-, δ-, ε- and ζ-keto acylimidazoles.

SCHEME 31.35

OH OH

O

OH O-Et

R1

O

– Cl–; – Im–

O

O

OH

1) +2e –, CTMS/TEA 2) TBAF/THF

Im

+e– +H+

+e– +H+

R2

O–Et –EtOH

O-Et R1

R2

OH O

O

R2

R1

R1

R2

V O

O R2 Reduction

O-Et R1

R2

O R2

R1 R1

VI

SCHEME 31.36

O

Electrochemical reduction of β-keto esters accompanied by the Tafel rearrangement.

conditions, the carbonyl is reduced and protonated and the radical is coupled with the electrophilic ester carbon, resulting in an intermediate cyclic pinacol that stabilizes to hydroxycyclopropanone. The latter, after the ring opening, yields an α-diketone, whose stepwise reduction finally results in the rearranged hydrocarbon (Scheme 31.36). The first part of the process—the pinacol formation— seems to be similar to the crossed pinacol coupling [51]. For this rearrangement, the β-carbonyl must be unconjugated (compound V). Esters like VI do not undergo this reaction and in acidic solution they are reduced to a secondary alcohol. When aromatic δ- and ε-keto esters are reduced, a similar intramolecular coupling proceeds, however, due to the alkoxy leaving group; the final product is an α-hydroxy ketone [130]. At the end of this section, an interesting example of intermolecular coupling of reduced diketones in nonaqueous DMF will be mentioned: Diaryl-1,2-diketones (e.g., benzil) are electrochemically reversibly reduced to the dianion that undergoes an intermolecular coupling with carbonimidoyl dichlorides under the formation of the corresponding enediol iminocarbonates (Scheme 31.37) [131]:

© 2016 by Taylor & Francis Group, LLC

1226

Organic Electrochemistry

Cl

+ O

+2e–

O–

N Cl O

–O

O

SCHEME 31.37 dichloride.

O

+2Cl–

N

Intermolecular coupling of reduced diaryl-1,2-diketones (e.g., benzil) with carbonimidoyl

G. REDUCTION MECHANISMS IN ANALYTICAL APPLICATIONS In Sections I.B through I.F, a number of model types of carbonyl compounds were selected and their characteristic electrochemical reduction mechanisms were presented and discussed. On the other hand, especially in electroanalysis, one can encounter important, biologically significant molecules containing one or more carbonyl groups, often with rather complicated structure and nonfitting unambiguously to one of the mentioned types. Nevertheless, even in these cases the reduction processes were investigated and the respective mechanisms were put forward. Several recent examples are presented in this section with the hope that the readers working with rather complex carbonylcontaining molecules can find here a similarity or an inspiration. The structures are depicted later. It is noteworthy that for trace analysis of reducible organic analytes, advanced polarographic and voltammetric methods with mercury electrodes are steadily considered to be very suitable. Due to their instantly renewable, absolutely smooth surface, they exhibit high reproducibility and sensitivity accompanied with simple preparation of samples (e.g., possibility of direct measurements of even suspensions or biological samples without separation or other pretreatment). In addition to this, small-size and low-cost instrumentation (often easily portable) are other advantages. For detection and determination of the agrochemical chlorophacinone in formulations, grains and vegetables a method based on differential pulse polarography (DPP) was developed. The suitable electrochemical response—DPP peak—was attributed to the simultaneous reduction of three carbonyl groups [132]. A spectrum of electroanalytical methods like dc-polarography (DCP), cyclic voltammetry (CV), DPP, controlled potential electrolysis (CPE), and millicoulometry was used for a study of the cathodic reduction of iprodione. A single four-electron signal interpreted as a simultaneous reduction of two carbonyl groups was obtained and was used for analytical purposes [133]. Valone is a pesticide containing a carbonyl group whose determination in water is important. CV and DPP study of fundamental electrochemical properties was performed and the signal at pH 4 was used for analytical application [134]. The antidepressant drug and smoking cessation aid bupropion hydrochloride (in fact its carbonyl group) is best reduced in slightly acidic medium on mercury electrodes. The electrochemical properties and kinetic parameters were evaluated [135]. Electrochemical and spectrophotometric study of the contraceptive drug norethisterone at mercury electrodes in aqueous solutions resulted not only in mechanistic considerations and determination of pKa, but particularly in its trace quantification in pharmaceutical formulations using stripping methods and reaching a detection limit 1.5 × 10−9 mol L−1 [136]. Reducibility and catalytic activity of lovastatin were investigated by polarographic and voltammetric methods. The main reduction wave at −1.49 V (vs. SCE) was ascribed to a two-electron and two-proton addition to the carbonyl group on the lactone ring. In the presence of a hydroperoxide, a strong catalytic reduction wave appears that can be used in trace detection of lovastatin (up to 8 × 10−9 mol L−1) in pharmaceuticals, urine, and blood serum [137].

© 2016 by Taylor & Francis Group, LLC

1227

Reduction of Aldehydes, Ketones, and Azomethines

The main cathodic response of captopril (a medical against hypertension) is the two-electron reduction of the carbonyl group at −0.9 V (vs. SCE). This signal was used for the development of a DPP detection of captopril in tablets [138]. A new polarographic (DC, AC, NPP, and DPP) method of daunomycin detection in pure form, vials, and in blood is presented, based on the cathodic peak of the quinone moiety. The analysis is complicated by a strong adsorption [139]. The mechanism for reduction of isatic acid was investigated by DCP and pulse polarography, linear sweep and CV, electrocapillary curves, chronocoulometry and bulk electrolysis with electrochemical and UV spectroscopic monitoring of the reaction. Three forms of isatic acid in aqueous solution were detected, depending on pH: diprotonated, monoprotonated, and nonprotonated. The reduction center is a carbonyl in all cases. In acidic media, adsorption was observed [140]. The mechanistic study of electroreduction of lichexanthone was performed in aprotic DMSO at stationary as well as rotating glassy carbon disk electrodes. The reduction proceeds in two one-electron steps to anion radical and dianion, the first is reversible. The slow follow-up chemical reaction involves opening of the 4-pyrone ring at the oxygen heteroatom and necessary protonation. The expected reduction of the carbonyl group does not take place [141]. An effect of substituents in position 7 on the reducibility of flavones was followed by NPP in DMF/aqueous buffer mixture and the corresponding kinetic parameters were evaluated using the Meites and Israel equation [142]. O

O

Cl O

Cl O N

Cl

O

O

O

O

captopril O

O CH3

H3C

H3C

O O valone

H 3C

lovastatin

O R

O

H

H O

O

O

OH

OH

CH3 OH

H

H O flavone

Me

N

OH

O

H3C

CH3

H3C

H 3C

H

O

O

O lichexanthone

© 2016 by Taylor & Francis Group, LLC

Me-O

CH3 OH

NH2 Me

O-Me

O

daunomycin

norethisterone

OH

OH

O

O

OH

O

+2e– +2H+ Me-O

O

HS

OH +2e– +2H+

H3C

O

HO

O

O

NH2 isatic acid

N

O iprodione

chlorophacinone HO O

COOH

H N

OH

H N

O-Me Cl

bupropion

1228

Organic Electrochemistry

II.

AZOMETHINE COMPOUNDS

A.

GENERAL MECHANISTIC CONSIDERATIONS

Azomethine compounds are molecules derived from carbonyls, where the oxygen is replaced by nitrogen. Since azomethines can be generally prepared by condensation of a carbonyl with an amine and the reverse hydrolysis yielding the starting substances is also possible, the azomethine group can be considered as a direct carbonyl analog with many common features with parent carbonyl compounds and simultaneously as a “krypto-” carbonyl, a precursor from which the carbonyl moiety can be released. Unlike carbonyls with oxygen being always the terminal atom, the azomethine bond need not to be only the terminal one, like in imines, but it is mostly incorporated in a chain or in a heterocycle. Therefore, the spectrum of various azomethine compounds is broader: the azomethine nitrogen can bear not only hydrogen, like in imines, but also carbon (Schiff bases, heterocycles), oxygen as hydroxyl (oximes) or as alkoxy substituent (O-alkyl oximes), and another nitrogen (hydrazones, azines), eventually other heteroatoms. Similar to carbonyl compounds, the azomethine bond is polarized with the electrophilic center attractive for nucleophiles (and electrons) at the carbon atom. However, the nitrogen with its nonbonding electron pair is prone to protonation, which affects significantly the reduction mechanism. Generally, two types of electroreduction mechanisms exist (Scheme 31.38): since the electron is the elementary nucleophile, the electrochemical reduction of neutral, unprotonated azomethine (in basic or aprotic media) starts by the electron transfer to the carbon atom under formation of a radical anion that either abstracts a proton from the solvent and is further reduced to saturation of the azomethine bond or that dimerizes [143–146]. On the other hand, in the case of protonated molecules, the first electron transfer is aimed at the nitrogen atom, causing first the splitting of the bond between two heteroatoms. In addition to this, a covalent addition of nucleophiles may play a significant role in reactivity and reducibility of azomethine compounds. The most frequent nucleophile is the hydroxide anion from water; hence, the discussion concerning the addition of nucleophiles will be devoted mainly to hydration. Addition of water to imines or Schiff bases derived from aliphatic or alicyclic carbonyls results in the corresponding carbinol amine, which is unstable and decomposes rapidly to the parent carbonyl and amine (hydrolysis of imines). In the case of benzylidene Schiff bases, the stability of the carbinol amine is higher due to the stabilizing aromatic system and depends on pH, benzaldehyde substituents, and structure of the amine [17]. However, carbinol amines are not reducible, therefore their formation and decomposition cannot be directly followed electrochemically and the proof of their existence must be done indirectly [147]. Covalent addition of water on the azomethine bond being a part of the unsubstituted N-heteroaromatic ring does not proceed without activation, which means without increased polarization of the C=N bond. This effect can be achieved either by substitution with an EWG in m-positions toward the

(–)

N

Y

R1 (+) R2

+e– Aprotic basic

–N

R1

Y

Dimerization protonation reduction of C=N

R2

(Y = O, N, C)

+H+

+ Y HN R1

SCHEME 31.38 respectively.

R2

+2e– Protic

NH R1

R2

+

_ Y

Protonation reduction of C=N (Y = O, N) hydrolysis

Electrochemical reduction mechanisms of unprotonated and protonated azomethine bonds,

© 2016 by Taylor & Francis Group, LLC

1229

Reduction of Aldehydes, Ketones, and Azomethines

heteroatom and/or by the second nitrogen in the ring (pyrimidine derivatives) or by another annealed aromatic ring [148] The first two conditions are well fulfilled, for example, in the case of 5-nitropyrimidine or its 2-methyl and 2-benzyl derivatives where covalent addition of water was found [149]. All azomethine groups (under conditions where the protonation steps may occur, that means, not only in aqueous solutions) are ultimately electrochemically reduced to amines. On the basis of the consumption of electrons during the electroreduction process, the compounds bearing azomethine bonds can be divided into three groups: Imines (>C=N–H) and Schiff bases (>C=N–R) derived from the reaction of a carbonyl with ammonia or primary amine are reduced by two electrons and two protons under saturation of the C=N double bond. Oximes (>C=N–OH or >C=N–OR) derived from the analogous reaction of carbonyl with hydroxylamine (NH2–OH) or alkoxyamine (NH2–OR), and hydrazones (>C=N–NR2) derived from hydrazine are reduced by four electrons and four protons since there are two reducible groups: the azomethine bond and the bond between the two heteroatoms N–Y (Y=O, N). Important factors influencing the reactivity of these azomethine compounds and stability of intermediates are (a) electron affinity of the double bond (π-electron system); (b) leaving ability of the group Y−. Azines (>C=N–N=CC=N–N=C (–O–R) R). Besides this fundamental classification, the individual reduction processes may differ (a) by the sequence of proton- and electron transfers; (b) by the order which group is reduced first: whether the C=N bond or the N–O (N–N) bond between two heteroatoms. These differences in reduction pathway depend naturally on the conditions (protic–aprotic solution and pH). Generally, azomethine compounds are reduced more easily than the parent carbonyls and antecedent protonation facilitates the reduction. Azomethine compounds—similar to the carbonyls—are frequently used, electrochemically investigated and their electrochemical reduction reviewed (e.g., References 9,150–153). In this part, a general overview is presented and supplemented by more recent achievements.

B.

DERIVATIVES OF AMMONIA (IMINES, SCHIFF BASES, IMINIUM CATIONS)

1. Protic Media Imines derived from aliphatic ketones are reducible only in alkaline solution; otherwise, they undergo a fast hydrolysis yielding the original carbonyl and ammonia. Preparative electroreduction of such imines to corresponding amines can be therefore performed only in a “one-pot” experiment, where the parent ketone reacts with the large excess of ammonia and the produced imine is simultaneously electrochemically reduced to the amine product (Scheme 31.39, left) especially at low temperature. The reduction potential of imine is less negative than that of the carbonyl. When a higher concentration of the ketone is used for electroreduction of the in situ generated imine, the product—amine—undergoes a “secondary” condensation with starting ketone yielding a new Schiff base (Scheme 31.39, right) [154]. Aromatic imines are more stable (and thus reducible) even in neutral and acidic media. Besides their formation from parent carbonyls, they are often encountered as reduction intermediates of oximes, hydrazones, and azines (see Sections II.C–II.E).

O

–H2O + NH3

N-H

+2e , +2H

NH2

+

O

–H2O N

SCHEME 31.39 “One-pot” condensation–reduction–condensation experiment resulting in an aliphatic Schiff base.

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1230

Organic Electrochemistry

H N

+e– +H+

R2 +H+

R1

+e–

H N

NH+

N R2

R2

R2

R1

SCHEME 31.40

H N

R2

Dim.

R1

R1

R1

R1

N H

R2

Electrochemical reduction of preprotonated aromatic Schiff bases in acidic media.

Schiff bases, especially the aromatic ones, are in water much more stable than imines. Their electrochemical reduction leads generally to a saturation of the azomethine double bond [155]. In acidic media, pre-protonated Schiff bases of aromatic carbonyls are reduced in two oneelectron waves that at higher pH merge into one (see Section I.A.1, Figure 31.1) [154,156]. The first wave corresponds (similar to the case of carbonyls) to radical formation, which may either dimerize (mechanism: CEDim) or be further reduced at the more negative potential in the second wave to saturation (CEEC) [143] (Scheme 31.40). 2. Aprotic Media In nonaqueous solvents, the azomethine compounds are more stable than in water. Their reduction occurs at more negative potentials because the nitrogen atom is not pre-protonated. Aliphatic ketimines and Schiff bases are typically reduced at rather highly negative potentials by a single two-electron process to a saturated C–N bond, following the ECEC mechanism. The primary imine radical anions are strong bases able to abstract proton(s) from the solvent and the resulting protonated neutral radical is generally reduced more easily than the parent Schiff base (Scheme 31.41, upper pathway) [144]. Aromatic Schiff bases, due to the stabilizing π-system, are reduced less negatively, their primary radical anions are less basic and more stable, hence their reduction proceeds in two one-electron waves according to an EEC mechanism (especially in DMF where activity of protons is lower than, e.g., in AN; Scheme 31.41, lower pathway). Besides, the radical anion may undergo a dimerization. Substituted iminium cations or protonated imines are reducible in a one-electron process in aprotic media giving rise to radicals that are relatively stable (especially those derived from (di) aryl ketones) and are able to dimerize [145]. The process is purely radical, without participation of protons and occurs at the originally carbonyl carbon [146].

+SH –S– +e– N R1

R2 R1

R2

+e–, +H+

R1 H N

Dimer

N– R2

H N

+2H+

+e– – – N

R2 R1

SCHEME 31.41

Electrochemical reduction of ketimines in nonaqueous media.

© 2016 by Taylor & Francis Group, LLC

R2 R1

1231

Reduction of Aldehydes, Ketones, and Azomethines

Alkyl

Alkyl

NH2

Alkyl

N–H

+2e–

+2e– +2H+

SCHEME 31.42

N–H

Alkyl

N–H +2H+

Preferential reduction of the aromatic ring in the case of 9-, 10- or 9,10-anthryl(di)imines.

When comparing electrochemical reduction (in DMF) of aromatic ketimines derived from benzene, naphthalene, and anthracene, the phenyl or 1-naphthyl ketimines are reduced in a standard way at the azomethine bond giving rise to the corresponding amine, while the 9-, 10-, or 9,10-anthryl (di) imines are reduced first at the aromatic ring under formation of the imine of 9,10-dihydroanthracene (Scheme 31.42). The chemical reduction showed the same difference in products; the experimental results were supported by quantum chemical calculations [157]. Aromatic imines, as direct analogs of quinones, where one or both C=O groups are replaced by N-aryl imino moieties (quinone mono- and diimines), are reduced by two electrons resulting in saturation of both C=N (C=O) bonds and aromatization of the ring [158]. More recently, a seminal study appeared, where based on a large series of newly synthesized quinone mono- and diimines, the structural influences on redox properties in aprotic as well as in protic media were systematically investigated and discussed [159]. It was found that the nature of the substituent on the nitrogen atom (N-phenyl-, N-phenylsulfonyl-, and N-benzoyl-) is the most significant indicator for nucleophilicity of radical intermediates and products. In aprotic media, the quinone monoimine VII is reduced in the standard way by two reversible one-electron steps followed by the abstraction of two protons from the solvent giving rise to p-aminophenol. By addition of proton donors of various strength, the nucleophilicity of the primary radical anion can be evaluated. In acidic solutions, a pre-protonation takes place, the iminium cation is reduced to the corresponding aminophenol, and this process is chemically reversible. The redox properties of VII were compared with analogous series of N-phenylsulfonyl- and N-benzoyl-imines and diimines with characteristic structures VIII–XII.

R4 R3

R1

R4

R2

R3 O

N

R5

R1

R4

N

VII

R1 R2

R3

R2 N

R5

R5

O

R4

R1

R3 O

R2

O

N O S O

S O N

R4

R1 R2

R3

N R4

R5

O

O

O

R1

R5 VIII

R5 IX

N

N O S O

R2

R3 N R5

X

R5 XI

R5 XII

Aromatic imines, with a good leaving group in the position 4 of the quinoid ring, have also similar character. In acetonitrile, the two-electron reduction causes splitting of the leaving group (in the presented case trichloromethyl) and the ring becomes aromatic simultaneously with the saturation of the azomethine bond. This reaction is significant mainly in organic electrosynthesis as a way to N-tolylation of aromatic amines (Scheme 31.43) [160].

© 2016 by Taylor & Francis Group, LLC

1232

Organic Electrochemistry +2e– R(H)

N

H N

R(H)

–C Cl3

C Cl3 CH3

H3C

SCHEME 31.43 Electrochemical reduction of aromatic imines with a good leaving group in position 4 of the quinoid ring.

In the case of newly synthesized aromatic Schiff bases combined with azobenzene XIII (exhibiting good antimicrobial activity), in DMF, the preferential reduction of the azomethine bond is reported [161]. O–Alk

H3C–O N N

N OH XIII

Electrochemical reduction of the enol–imine compound XIV in DMF proceeds in two oneelectron quasireversible steps (to radical anion and dianion), according to expectations. The reversibility is suppressed by the presence of two hydroxylic hydrogen atoms due to the intramolecular hydrogen bonding and autoprotonation reaction [162].

H O HO

H3C CH3 CH3

N

XIV

When aromatic Schiff bases (fluorinated benzylidene anilines) are reduced in aprotic conditions in the presence of CO2, the primary radical anion reacts with the carbon dioxide under formation of fluorine-containing amino acids and electrochemically activated insertion of carbon dioxide is going on (Scheme 31.44). The originally reversible first wave becomes irreversible and increases, whereas the second reduction step (to the dianion) disappears [163,164]. In the following, more recent study, the crucial influence of ion-pairing (various solvents, various indifferent electrolytes) on effectiveness is discussed [165]. F

F

F

F

– N

+e–

N–

CO2

N OOC

R

R

H N

+e–, +2H+ OOC

R

R

SCHEME 31.44 Electrochemical reduction of fluorinated benzylidene anilines in aprotic conditions in the presence of CO2.

© 2016 by Taylor & Francis Group, LLC

1233

Reduction of Aldehydes, Ketones, and Azomethines

+2e– +CTMS

R

+

R

O

H3C

R H3C

Im

_ N

CH3

–TMS

O

R

HN

N TMS

TMS

SCHEME 31.45

O– Im

–Cl–

N

R –Im–

R

R

R

Coupling of a reduced diarylimine (Schiff base) with a ketone.

Another synthetic application of electrochemical imine reduction in aprotic media is based on coupling of a reduced diarylimine (Schiff base) with a ketone (acylimidazole is very suitable since imidazole is a good leaving group) in the presence of chlortrimethylsilane (CTMS). In this process, α-amino-αaryl ketones (C-acylated imines) are formed enabling synthesis of, for example, α-aryl glycine (Scheme 31.45, see Scheme 31.17) [166]. This type of reaction can be used also intramolecularly [167].

C.

DERIVATIVES OF HYDROXYLAMINE (OXIMES, O-ALKYLATED OXIMES, N-ALKYLATED OXIMES–NITRONES)

1. Protic Media Oximes are generally reduced to amines by four electrons. There were two theoretical possibilities of the first reduction step: saturation of the C=N bond or splitting of the N–O bond. Lund suggested already in the 1950s that in acidic or neutral aqueous media the first two-electron reduction proceeds at the N–O bond under formation of imine since the hypothetic corresponding hydroxylamine is nonreducible under the given conditions [154]. The intermediate imine is then reduced in the second two-electron step to the final product— amine, or, in the case of (di)alkylimine and/or in slightly acidic pH, the imine is simultaneously hydrolyzed to the original carbonyl compound and ammonia. Which of these processes prevails (electroreduction vs. hydrolysis) depends on the imine structure (aromatic imines are more stable intermediates than aliphatic ones) and on conditions, particularly pH. The reduction of the intermediate imine can occur either at less negative potentials than the starting oxime and a single four-electron process is observed, or at more negative potentials and two 2-electron waves are obtained (this is possible in strongly acidic media). In the latter case, the imine is relatively stable to be isolated and identified. This was achieved in the case of a cyclic oxime 4-(4′-methoxyphenyl)-2, 3-benzoxazin-1-one (Scheme 31.46) and 2,4-dihydroxybenzophenone oxime [168], when the existence of both intermediate imines was experimentally proven. Later the presence of imines during the reduction of a series of p-substituted aryl aldoximes and ketoximes was confirmed for those derivatives when the p-substituent was an electron-donating group [169]. O-CH3

+2e–, +2H+

+H+ N

NH+

O

O

O

O-CH3

O-CH3

O

NH2+ OH O

SCHEME 31.46 Electrochemical reduction of a cyclic oxime in strongly acidic media yielding a stable protonated imine.

© 2016 by Taylor & Francis Group, LLC

1234

Organic Electrochemistry

In acidic and neutral solutions, oximes—like other azomethines—exist in the protonated form, hence depending on pH, the species being actually reduced differ. This feature is evident from (a) the total number of observed reduction processes (waves) within the total pH range (at lower pH the protonated positively charged molecule is reduced at less negative potentials than the neutral species in basic solution), and (b) from the linear dependence of the reduction potential on pH, pointing to a pre-protonation reaction in the diffusion layer. There is a general rule (see Section II.A): when the molecule itself prevailing at the given pH in the bulk of solution is reduced, no antecedent proton transfer takes place and the reduction potential is independent of pH. Since the protonated form is always reduced more easily, at slightly higher pH when the prevailing bulk form is already nonreducible at the given potential, the molecules transported from the bulk are continuously protonated in the diffusion layer and immediately reduced. The depletion of the pre-protonated form shifts the equilibrium to its side and therefore the full reduction may occur even 3–4 pH units above the pKa. Since the supply with protonated form with increasing pH is more difficult, the reduction potential is linearly shifted to more negative values. The slope of this dependence indicates among others the number of protons involved. At the oxime moiety, there are two sites of possible protonation (N and O), therefore both should be considered. Based on the respective pKa, in strong acids a diprotonated form prevails in the bulk, in slightly acidic media monoprotonated species dominate. In neutral to basic solution, the uncharged original form of the molecule is typical, which in very alkaline environments dissociates to the irreducible oximate anion. The protonated oximes in acidic–neutral media are thus always reduced by four electrons to amines, with imines as intermediates. In basic solutions, however, (and in aprotic solvents—see Section II.C.2) the situation is different and the most electrophilic site on the unprotonated molecule is the oxime carbon. Therefore, in the case of reduction of benzaldoxime in alkaline media, only two-electron reduction occurs under the formation of benzyl hydroxylamine [170]. In a more recent study of electroreduction mechanism of a series of substituted benzaldehydeand acetophenone oximes [171], a diprotonation mechanism was suggested based on the shape (slope) analysis of the E1/2 versus pH and ilim versus pH dependence. In the diprotonated structure >C=NH+–OH2+, the positive charges at both heteroatoms focus the electron uptake to this site and the cleavage of the N–O bond is facilitated by the good leaving group (OH2). This process occurring at pH between 3 and 8 proceeds in close vicinity of the electrode. It is adsorption assisted and analogous to the reduction mechanism of nitrones (N-alkylated oximes), where antecedent protonation of the already positively charged molecule takes place. This represents another example of the cleavage of the N–O bond with positive charges on both the nitrogen and the adjacent oxygen atoms [172]. In the case of trifluoroacetophenone oxime, a strong hydration was found due to the strong electron-withdrawing effect of –CF3. Between pH 2 and 8, the reduction limiting current decreases to about 2% of the original value recorded at pH < 2, pointing to the nearly exclusive presence of carbinolamine, a species assumed to be the first unstable intermediate in reaction of carbonyl compounds with amines (Scheme 31.47) [173]. The four-electron reduction pattern yielding the corresponding amine operates also in pyridine-4-aldoxime [174] as well as in the series of methyl hetaryl (pyridin-2-yl, pyridazin-3-yl,

OH N

HN

OH +H2O

CF3

SCHEME 31.47

Hydration of trifluoroacetophenone oxime.

© 2016 by Taylor & Francis Group, LLC

OH CF3

1235

Reduction of Aldehydes, Ketones, and Azomethines

pyrimidin-2-yl, pyrimidin-4-yl, pyrazinyl) ketoximes [175]. Due to the presence of more heteroatoms, the acidobasic properties are more complex (pre-protonation, shift of reduction potentials, formation of intramolecular hydrogen bonds etc.). Nevertheless, depending on pH, only three different acidobasic forms of hetaryl oximes (di-, mono-, and unprotonated) could be identified during reduction of all studied derivatives. Within the whole pH scale, three 4-electron processes were observed that replace each other with increasing pH [175]. In the case of the pyridine-4-aldoxime, at potentials close to the electrolyte discharge, the second reduction process occurs assigned to the reductive splitting of the amine [174]. Changes of the δE1/2/δpH and δilim /δpH dependences over the whole pH region were presented and discussed in the context of protonation state, kinetics and (α. n) values. Imines (even unstable in acidic solutions) as intermediates are expected at all pH. In the diazine ketoximes, the second ring nitrogen (according to its position) facilitates reduction of the oxime group in comparison to methylpyridin-2-yl ketoxime. Reducibility of the oximes correlates with the enhancement of their hydrolytic activity towards organic esters (4-nitrophenyl acetate served as a model). The reduction of the heteroaromatic ring in the diazine ketoximes proceeds at more negative potentials than the reduction of the oxime group [175]. Nevertheless, one should keep in mind that pyridine-4-aldoxime XV can theoretically exist in two tautomers: besides the classical oxime formula, the quinoid tautomer is also reported [176], hence, the acidobasic properties as well as the reactivity of this compound can be more complicated. In the electrochemical literature, however, no such discussion has been found. O

OH N

N

N

N H

XV

Derivatives alkylated on the oxygen atom (O-alkyloximes) are another type of oximes. Their reduction mechanism in buffered solutions was found to be analogous to that of other oximes. In the case of the antibiotic cefetamet, a single four-electron process in acidic media is replaced by two two-electron waves at pH 5–9 corresponding to a gradual reduction to the imine and amine, respectively. The reason for the separation of potentials of the two processes was explained on the basis of combination–superposition of two acid–base equilibria causing pre-protonation of the starting oxime and intermediate imine, respectively [177]. O

HO

NH2

O

H3C

N

O

N S

S H

N H

N O CH3

cefetamet

The oxime reduction can be complicated by the spontaneous hydrolysis of the imine or in the presence of alcohol by its covalent addition to the intermediate imine. As a result, a decrease of the limiting current is observed. It was demonstrated that at a given concentration of the alcohol, the drop of current increases in the sequence MeOH < EtOH < i-PrOH < tert-BuOH, following the increasing nucleophilicity of the alcohol [178].

© 2016 by Taylor & Francis Group, LLC

1236

Organic Electrochemistry

2. Aprotic Media Electrochemical reduction of benzaldehyde oximes in DMF in presence of a proton donor (e.g., phenol) proceeds as a totally four-electron process, resulting in the respective amine. In carefully dried DMF, the presence of the intermediate anion radical PhCH=N− • after the rupture of the N–O bond was suggested. Among the products, N,N′-dibenzylhydrazine and benzonitrile were detected [179]. In the only, up to now published systematic study of the reduction mechanism of oximes in dry DMF, six types of oximes were discussed [180]. In electroreduction of benzaldoximes, their O-alkyl and O-acyl derivatives, the formation of benzonitrile as a product was confirmed. (This finding is rather surprising since the conversion of benzaldoxime to benzonitrile is not a reduction.) Two characteristic features were experimentally observed: (a) the CV curve starts with a small irreversible reduction peak, which is “interrupted” by a well-developed reversible couple of peaks of the stable product, identified as benzonitrile; (b) the charge consumption during preparative electrolysis resulting in 60–80% yield of benzonitrile corresponds to n = 0.4–0.6 only. The formation of nitrile from benzaldoxime was explained by a combination of electroreduction and electrocatalytic processes (Scheme 31.48) initiated by a one-electron reduction of the oxime followed by the splitting of the N–O bond under the formation of the imine radical and a hydroxide anion as an electrogenerated base (EGB). The OH− attacks the parent oxime resulting in elimination of water giving rise to benzonitrile. The alternative pathway is the deprotonation of the imine radical and the formed nitrile radical anion transfers the electron to the parent oxime or back to the electrode. The reduction of O-alkyl and O-acyl derivatives follows an analogous pattern and differs only in the base being split-off. In acetophenone, fluorenone, and O-acyl fluorenone oximes, benzonitrile cannot be formed due to the lacking hydrogen on the oxime carbon. Their preparative reduction in aprotic DMF yields the corresponding imine that is hydrolyzed by the present base and/or during the work-up to acetophenone and fluorenone, respectively [180]. The electroreduction behavior observed for pyridine-4-aldoxime in DMF differs from that described earlier in this section for benzaldoxime: in the cyclic voltammograms, no reversible peak appeared that could be ascribed to the reduction of the pyridine-4-nitrile; on the other hand, the final reduction product was identified as the amine derivative. The expected mechanism starts as in the previous case by the uptake of the first electron and splitting the N–O bond to imine radical and hydroxide anion. The latter, however, deprotonates the starting pyridine-4-aldoxime giving rise to the oxime anion, which is reduced more negatively and abstracting protons from the solvent yields the corresponding amine [181]. From the electrosynthetic point of view, an electroreductive intramolecular coupling of O-methyloximes with aliphatic cyclic imides in 2-propanol afforded five-, six-, and seven-membered cyclized products. These reactions provide a useful method to synthesize azabicyclo compounds (Scheme 31.49) [182].

— N–OH + e PhCH —

–

OH– + PhCH — — N–OH — N° + OH– PhCH —

– PhCH — — N° + OH

PhCH — N–OH–° H2O PhC

PhC

N–°

+

PhCH — — N–OH

PhC

N + e–

PhC

N–°

+

PhC

N

+ OH–

+ H2O PhC

N

+ PhCH — N–OH–°

N–°

SCHEME 31.48 Proposed mechanism of the formation of nitrile from benzaldoxime.

© 2016 by Taylor & Francis Group, LLC

1237

Reduction of Aldehydes, Ketones, and Azomethines

O

H N O–Me

N

n

R

O

Electroreduction N n O

N

O–Me

Work-up

H

H O–Me N

N

n

i-PrOH O

O

n = 1–3

SCHEME 31.49 Electroreductive intramolecular coupling of O-methyl oximes with aliphatic cyclic imides yielding azabicyclo compounds.

3. Mono- and Dioximes of α-Diketones The presence of the electron-withdrawing carbonyl group in α-position with respect to an oxime enhances the reducibility of the oxime function. A similar rule is valid also for O-substituted oximes = oxime alkylethers, and monohydrazones of α-diketones. As a result, the oxime is reducible even at pH 10.2 when the oxime is in its anionic form. The reduction process in aqueous media is going on in a standard way, like in isolated oximes: a four-electron reduction results in an α-amino carbonyl compound, which is at more negative potentials further reduced and ammonia is released [183,184]. It has been proposed that in aromatic species also the enaminol might be the possible alternative intermediate [185,186]. This problem was later studied on the example of aryl-alkyl diketones and the regioselectivity of the reduction was demonstrated using the comparison of behavior of two isomeric monooximes of phenyl-methyl diketone. At pH lower than 5 in the time scale of seconds (under conditions of d.c. polarography and CV), the 1-phenyl-1,2-propanedione-1-oxime (XVI) is reduced in a standard way by four electrons to 1-phenyl-1-aminopropan-2-one. The isomeric 1-phenyl-1,2-propanedione2-oxime (XVII), however, does not yield the expected reducible α-aminoketone but a nonreducible olefin derivative (oxidizable to ketoimine) [187]. In the time scale of tens of minutes (during controlled potential electrolysis), the olefinic enaminol undergoes a slow conversion to the 1-phenyl-2aminopropan-1-one as the main, further reducible product. The optimized geometry showed that the protonated imine of XVI is in cisoid nonplanar conformation and the LUMO is unsymmetrical, localized mainly at the imine moiety. Hence, the imine group is rather isolated and the delocalization with the carbonyl is limited. On the other hand, the protonated imine of XVII exists in transoid, nearly planar conformation with the LUMO symmetrically spread over the two carbonyl functions. This molecule (including the LUMO) closely resembles benzil that is reduced to the stable and further nonreducible enediol. This comparison suggests the role of the molecule conformation in the reduction mechanism of monooximes of α-diketone (Scheme 31.50) [188]. The vic-dioxime is reduced in acidic-neutral media in a six-electron wave under formation of an ene-diamine further reducible at more negative potentials. The last reduction step is again H3C O N

OH

+2e– +2H+

H3C O

–H2O

+2e– +2H+

H3C O

NH

NH2

“cisoid”

XVI OH N CH3 O

+2e– +2H+ –H2O

XVII

HN

O “transoid”

+2e– CH3 +2H+

H2N

H2N CH3 OH

CH3 O

SCHEME 31.50 Role of the molecule conformation in the reduction mechanism of two isomeric monooximes of α-diketone.

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Organic Electrochemistry

N

N–OH

+6e– +6H+ –2H2O

NH2

NH

Slow

NH2

NH2

NH2

N–OH

+2e– +2H+

O

Hydrolysis

O

–NH3

+2e– +2H+ – NH2 +2e + +2H

SCHEME 31.51

NH2

NH

Slow

O

Hydrolysis

Six-electron electrochemical reduction of the vic-dioxime in acidic–neutral media.

analogous to that of benzil. The stability of the ene-diamine, however, is limited, and it is slowly transformed to the amine-imine undergoing hydrolysis to the α-aminoketone. From the latter, ammonia is reductively split-off at even more negative potentials. Using controlled-potential electrolysis at the first six-electron wave, the α-aminoketone was identified as the main product (Scheme 31.51). Depending on the pH and structure (aromatic–aliphatic, aldimines–ketoimines), the protonation and concurrent reactions (tautomeric changes, hydrolysis) proceed in various proportions [186,189].

D.

DERIVATIVES OF HYDRAzINE (HYDRAzONES)

1. Protic Media The reduction of some hydrazones may be influenced by tautomeric changes to “azo-” or “enehydrazine” forms when the structure permits the hydrogen transfer (Scheme 31.52). The hydrazone form is dominant in acidic and neutral media, the other tautomers are expected in basic solution [190,191]. This can be the reason for the low reducibility of hydrazones in basic media. Phenylhydrazones of aromatic carbonyl compounds are reduced in acidic media in a four-electron process: first, the protonated form of the central N–N bond is split by two electrons, to an aniline derivative and a corresponding aromatic imine, which is subsequently reduced by other two electrons under saturation of the azomethine bond (Scheme 31.53) [154]. When the amine nitrogen is bearing an EWG, and, especially in basic media, the antecedent protonation becomes difficult. Then the electron uptake occurs first and a two-electron reduction takes place yielding a hydrazine derivative (Scheme 31.54). The same reaction pattern is followed by acylated hydrazones [192]. Phenylhydrazones derived from aliphatic carbonyl compounds are not reducible in aqueous solutions, only their trialkylhydrazonium salts [9]. Hydrazones derived from aromatic aldehydes and ketones are reduced at pH 2 to about 8 in a four-electron step. The observed linear pH-dependence of half-wave potentials of N,N,N-trialkylhydrazonium ions indicates a protonation of the azomethine nitrogen prior to the first electron uptake. The reduced species is thus bearing two positive charges on adjacent nitrogen atoms.

N

R1

R3 N

N

R2 azo-

SCHEME 31.52

R3 N

R1

R3 H

R2 hydrazone

Tautomeric equilibria of some hydrazones.

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HN

R1

N

H

R2 ene-hydrazine

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Reduction of Aldehydes, Ketones, and Azomethines

H+ N

+H+ N

Acidic soln.

N

H N

N

N R2

R1

R2

R1

R2

R1

+e–

+H+ +e–

+2e–, 2H+ NH

NH2 R1

R1

SCHEME 31.53

R1

R1

R2

Four-electron electrochemical reduction of phenylhydrazones in acidic media.

+e–, +H+ N

H + N H N

– H N R2

N EWG

H N

Basic soln.

N EWG

R1

+e–

H – N R1

SCHEME 31.54

+H+ N EWG

H H N R1

N EWG

Two-electron electrochemical reduction of phenylhydrazones in basic media.

From the evaluation of other δE1/2/δpH dependencies in acidic-neutral media plotted for a series of hydrazones and oximes follows that their electrochemical behavior is consistent with the formation of a dicationic form, present as a reducible intermediate at the surface of the mercury electrode [193,194]. The proven presence of imines as intermediates represents a confirmation of the reduction pathway where the initial two-electron reductive cleavage of the N–N bond takes place [195]. In the case of hydrazones derived from α,β-unsaturated carbonyl compounds (e.g., of carvone or cinnamaldehyde), in alkaline solutions the α,β-double bond is reduced first by two electrons under formation of a saturated hydrazone [9,154]. 2. Aprotic Media For the reduction mechanism of aromatic hydrazones ArR1C=N–NR2R3 in aprotic media, the principal question is whether one of the R-substituents is hydrogen. When R1, R2 and R3 are carbon substituents (structure type XVIII), the reduction starts with a reversible one-electron step under formation of a stable radical anion, especially in phenylhydrazones (Scheme 31.55, upper part). The presence of protons results in splitting of the N–N bond to imine radical and amine anion and their reduction/protonation yielding primary and secondary amines. In this sense, the process is analogous to the reduction of aromatic vinyl halides [196] or aryl halides [197,198]. When R1 is a hydrogen (derivative of benzaldehyde), the N-dialkyl hydrazone is deprotonated by an EGB, the amine anion (as a leaving group) is eliminated and a nitrile is formed (Scheme 31.55, lower part) [199–201]. When R2 or R3 is a hydrogen, its acidity is high enough to protonate the primary radical anion. The reduction becomes irreversible [202] due to the autoprotonation (father-son) mechanism [203] and the C=N bond is reduced [204].

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Organic Electrochemistry

XVIII

N N

–N

N

R2 +e–

R1

R3

R3

–

R3

XVIII

+

HN R2

NH2 Protonation reduction

R1

+

R2

R1

R3 N N

N R2 R3

H

SCHEME 31.55

+

EGB–

+

–N

+

H-EGB

R2

Electrochemical reduction mechanism of aromatic hydrazones in aprotic media.

E. AzINES (CYCLIC, ACYCLIC) 1. Protic and Aprotic Media Benzalazine (XIX) has two azomethine bonds; therefore, compared to hydrazones, two more electrons (and protons) are needed for its electrochemical reduction. A six-electron reduction occurs in acidic aqueous solution giving rise to two molecules of benzylamine. In basic solution, the reduction yields benzaldehyde benzylhydrazone after consumption of two electrons and two protons. The latter is further reducible (by four electrons to benzylamine) only in acidic media [9]. The question remains, what is the initial reaction in acidic solution, or, in other words, whether benzaldehyde benzylhydrazone is the first intermediate even in acidic media or if the first reduction step causes splitting of the N–N single bond. This discussion was reopened by the study of 2-benzoylpyridine azine [205] where the pathway involving azo- and hydrazine derivatives was suggested. Based on the analogy concerning the antecedent protonation of other azomethine compounds (oximes, hydrazones) preferentially at the two adjacent heteroatoms, the most probable general reduction mechanism of azines in acidic and neutral media starts with (di)protonation of the two heteroatoms (stabilized at the electrode surface), followed by a two-electron reductive splitting of the N–N bond (Scheme 31.56—see also Sections II.C and II.D on oximes and hydrazones). This mechanism was supported in the study of electrochemical reduction of unsymmetrical azines derived from aromatic and aliphatic carbonyls in aqueous acidic or neutral solutions [206]. It was concluded that the central single N–N bond is split into two imines first. The following reduction of the azomethine bond occurs, however, only at the aromatic imine (due to its higher stability) under formation of the corresponding amine. The aliphatic imine, on the other hand, undergoes fast hydrolysis, resulting in the parent dialkylketone and ammonia. Hence, in the latter case, only four electrons were consumed for the total reductive degradation of the studied aryl-alkyl azine. On the other hand, the preferential reduction of one of the azomethine bonds in alkaline solution is consistent with the reduction of the same type of azines in non-aqueous solvents (acetonitrile—AN, DMF). Under strictly aprotic conditions, benzalazine is reduced in two one-electron reversible steps under the formation of radical anion and dianion where both electron transfers are aimed to the unprotonated azomethine bonds [207]. The stability of the radical intermediate allows to record its EPR spectrum due to the presence of the stabilizing aromatic system. The primary radical anion of fluorenone azine is thus stable in DMF for months. In the presence of a proton donor (e.g., hexafluoro-2-propanol), an analogous two-electron reduction like in alkaline solution occurs where one azomethine bond is saturated giving rise to fluorenone fluorenylhydrazone [208,209].

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Reduction of Aldehydes, Ketones, and Azomethines

N

H+ N

+2H+ N

Acidic soln.

+2e–

HN 2

N+ H

XIX

+4e_, +4H+

HN

2

2

H2N

– N

+e_

N

Basic soln.

N

XIX

– N

SCHEME 31.56

XIX

XIX

Aprotic

N H

2–

+e–

+e– N

+H+, +e–, +H+

…

Electrochemical reduction mechanism of azines.

It is evident that the presence of protons decides about the reduction mechanism of azines. The reduction pathways of benzalazine-type compounds under various conditions are summarized in Scheme 31.56. Though the azine grouping >C=N–N=C< seems to be an electron delocalized conjugated π-system analogous to 1,3-butadiene, it was found that in fact the single N–N bond as well as the C=N double bonds are isolated and noncommunicating. The latter feature was detected in cyclic azines (1,2,4-triazine derivatives XX and XXI) [210–212]. Unlike the acyclic azines, the first step of electrochemical reduction of these molecules in aqueous media involved the saturation of the 1,6-C=N bond (proven by preparative electrolysis), at more negative potentials the 2,3-azomethine bond was reduced without a ring opening. As a model compound, the 2,3-dihydroderivative was chemically independently prepared and its 1,6-azomethine bond was electrochemically reduced. Surprisingly, the reduction potential of the 1,6-C=N bond of these two compounds were identical. That means that the reduction center at the 1,6-azomethine bond is not influenced by the presence or absence of the neighbor double bond being in structural conjugation. Hence, the N–N bond should block the electron delocalization in this heterocyclic azine [213]. NH2 4

O

NH2

N

Me

5

O

N

t-Bu

N

N

N Ph

S-Me

3 6

N

2 1

XX

XXI

To prove this phenomenon also in acyclic azines, the extent of electron delocalization in symmetrical and unsymmetrical p-substituted benzalazines [214] and in a series of analogous aryl-substituted acetophenone azines [207] was systematically studied using detailed interpretation of electroreduction data of the mentioned compounds in AN and DMF. The results proved the existence of a reduction system containing two localized, noncommunicating redox centers. The N–N

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1242

Organic Electrochemistry O-R

O-R +e– (Aprotic)

O-R

R3

R1

N

+2H+, +e– Solv-H

R3

N

H H N R1 R2

R2 N

R3

N–

R1

N R2

R3

O-R

H

O-R +H+ (Protic)

N

N+

R1

+2e–, +H+

R2

SCHEME 31.57

R3

NH

+

R1

HN R2

Electrochemical reduction mechanism of hydrazonates.

bond should have thus a single bond character. The comparison of x-ray structural data of various conjugated analogues of 1,3-butadiene revealed that whereas the bond order of the central C–C bond in 1,3-butadiene is about 1.4, the bond order of the N–N bond in acyclic azines is below unity (0.9). This surprisingly long N–N bond in azines is well consistent with the electrochemically observed blocking of electron delocalization along this grouping [214,215]. This problem was treated theoretically and in the solid state by Glaser [216,217]. 2. Hydrazonates Hydrazonates are formally azines where the –O–R group represents one substituent on carbon atom. Hydrazonates are reduced in aprotic media on a mercury electrode in two steps; the first one is a reversible one-electron process, but affected by a rather fast homogeneous follow-up reaction, since the ratio Ipa /Ipc is less than one, the second reduction step is irreversible. When an EWG group is present as the substituents R1 or R2, a third reduction step appears at more negative potential. Controlled-potential electrolysis in aprotic conditions at the potential of the first wave lead to saturation of a C=N double bond (Scheme 31.57, upper part). When a proton donor (phenol) is gradually added, a new, most positive and irreversible reduction peak appears and increases on expense of the originally first one. This effect points to a pre-protonation of one nitrogen atom leading to a change of mechanism where the protonated nitrogen atom is the reduction center. The preparative electrolysis in presence of a proton donor then results in the splitting of the N–N bond (Scheme 31.57, lower part). As a product, only imine (R1 = Me, R2 = Aryl) could be isolated and identified, whereas the iminoether undergoes fast hydrolysis during the workup [218]. Hence, the reduction behavior of hydrazonates is in principle analogous to that of azines.

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63. 64. 65. 66. 67. 68.

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164. Bondarenko, V. N.; Titov, V. E.; Koshechko, V. G.; Pokhodenko, V. D. Theor. Exp. Chem. 2008, 44, 271–277. 165. Titov, V. E.; Bondarenko, V. N.; Koshechko, V. G.; Pokhodenko, V. D. Theor. Exp. Chem. 2010, 46, 8–13. 166. Kise, N.; Morimoto, S., Tetrahedron 2008, 64, 1765–1771. 167. Kise, N.; Ohya, K.; Arimoto, K.; Yamashita, Y.; Hirano, Y.; Ono, T.; Ueda, N., J. Org. Chem. 2004, 69, 7710–7719. 168. Lund, H.; Acta Chem Scand. 1964, 18, 563–565. 169. Celik, H.; Ludvík, J.; Zuman, P. Electrochim.Acta 2006, 51, 5845–5852. 170. Lund, H.; Tetrahedron Lett. 1968, 9, 3651–3654. 171. Celik, H.; Ekmekci, G.; Ludvík, J.; Pícha, J.; Zuman, P. J. Phys. Chem. B 2006, 110, 6785–6796. 172. Zuman, P.; Exner, O. Collect. Czech. Chem. Commun. 1965, 30, 1832–1852. 173. Celik, H.; Ludvík, J.; Zuman, P. Electrochim. Acta 2007, 52, 1990–2000. 174. Roman, A. J.; Sevilla, J. M.; Pineda, T.; Blazquez, M. J. Electroanal. Chem. 1996, 410, 15–20. 175. Cibulka, R.; Liška, F.; Ludvík, J. Collect. Czech. Chem. Commun. 2000, 65, 1630–1642. 176. http://pubchem.ncbi.nlm.nih.gov/compound/Pyridine-4-aldoxime, last accessed February 2015. 177. Kapetanovic, V.; Aleksic, M.; Zuman, P. J. Electroanal. Chem. 2001, 507, 263–269. 178. Aleksic, M.; Kapetanovic, V.; Zuman, P. Collect. Czech. Chem. Commun. 2001, 66, 1005–1010. 179. Vagina, G. A.; Troepol’skaya, T. V.; Kitaev, Yu. P. Bull. Acad. Sci. USSR, Div. Chem. Sci. (Engl.Transl.) 1983, 32, 2237–2241. (Translated from Izv. Akad. Nauk SSSR, Ser. Khim. 1983, 2488–2493.) 180. Soucaze-Guillous, B.; Lund, H.; Acta Chem. Scand. 1998, 52, 417–424. 181. Roman, A. J.; Sevilla, J. M.; Pineda, T.; Blazquez, M. J. Electroanal. Chem. 2000, 485, 1–6. 182. Kise, N.; Fukazawa, K.; Sakurai, T. Tetrahedron Lett. 2010, 51, 5767–5770. 183. Andruzzi, R.; Cardinali, M. E.; Trazza, A. Electrochim. Acta 1972, 17, 1524–1528. 184. Cardinali, M. E.; Carelli, I.; Andruzzi, R. J. Electroanal. Chem. Interfacial Electrochem. 1973, 47, 335–342. 185. Armand, J.; Boulares, L.; Pinson, J.; Souchay, P. Bull. Soc. Chim. Fr. 1971, 1918–1919. 186. Armand, J.; Bassinet, P.; Boulares, L. C. R. Acad. Sc. Paris Ser. C 1973, 277, 695–698. 187. Celik, H.; Ludvik, J.; Zuman, P. Electrochem. Commun. 2006, 8, 1749–1752. 188. Ludvík, J.; unpublished results, to be submited. 189. Cardinali, M. E.; Carelli, I.; Trazza, A. J. Electroanal. Chem. Interfacial Electrochem. 1973, 48, 277–283. 190. Simon, H.; Modlenhauser, W. Chem. Ber. 1967, 100, 1949–1960. 191. Ioffe, B. V.; Stopskij, V. S. Tetrahedron Lett. 1968, 9, 1333–1338. 192. El Baradie, H. Y. F.; Ghoneim, M. M.; Issa, R. M.; Madkour, L. M. Bull. Electrochem. 1987, 3, 23–27. 193. Baymak, M. S.; Celik, H.; Ludvík, J.; Lund, H.; Zuman, P. Tetrahedron Lett. 2004, 45, 5113–5115. 194. Baymak, M. S.; Celik, H.; Lund, H.; Zuman, P., J. Electroanal. Chem. 2006, 589, 7–14. 195. Baymak, M. S.; Celik, H.; Lund, H.; Zuman, P. J. Electroanal. Chem. 2005, 581, 284–293. 196. Gatti, N.; Pedersen S. U.; Lund, H. Acta Chem. Scand. 1988, 42B, 11–22. 197. Savéant, J.-M. Acc. Chem. Res. 1993, 26, 455–461. 198. Savéant, J.-M. J. Phys. Chem. 1994, 98, 3716–3724. 199. Soucase-Guillous, B.; Lund, H.; J. Electroanal. Chem. 1997, 423, 109–114. 200. Kargin, Yu. M.; Latypova, V. Z.; Kitaeva, M. Yu.; Vafina, A. A.; Zaripova R. M.; Ilyasov, A. V. Izv. Akad. Nauk SSSR Ser. Khim. 1984, 2206; 2410. 201. Kargin, Yu. M.; Kitaeva, M. Yu.; Latypova, V. Z.; Zaripova, R. M.; Ilyasov, A. V. Izv. Akad. Nauk SSSR Ser. Khim. 1988, 510–514; 607. 202. Gudeika, D.; Lygaitis, R.; Mimaité, V.; Grazulevicius, J. V.; Jankauskas, V.; Lapkowski, M.; Data, P. Dyes Pigm. 2011, 91, 13–19. 203. Maran, F.; Roffia, S.; Severin, M. G.; Vianello, E. Electrochim. Acta 1990, 35, 81–88. 204. Sethukumar, A.; Arul Prakasam, B. J. Mol. Struct. 2010, 963, 250–257. 205. Gomez Nieto, M. A.; Luque de Castro, M. D.; Valcarel, M. Electrochim. Acta 1983, 28, 1725–1732. 206. Fuhlendorff, R.; Lund, H. Acta Chem. Scand. 1988, 42 B, 52–54. 207. Sauro, V. A.; Workentin M. S. J. Org. Chem. 2001, 66, 831–838. 208. Kitaev, Yu. P.; Ivanova, V. K.; Mukhtarov, A. S.; Orlova, L. N.; Ladygin, A. Izv. Akad. Nauk SSSR Ser. Khim. 1974, 64–66, 72. 209. Trieve, F. M.; Hawley, M. D. J. Electroanal. Chem. 1981, 125, 421–435. 210. Riedl, F.; Ludvík, J.; Liška, F.; Zuman, P. J. Heterocycl. Chem. 1996, 33, 2063–2064. 211. Ludvík, J.; Riedl, F.; Liška, F.; Zuman, P. Electroanalysis 1998, 10, 869–876. 212. Ludvík, J.; Riedl, F.; Liška, F.; Zuman, P. J. Electroanal. Chem. 1998, 457, 177–190. 213. Zuman, P.; Ludvík, J. Tetrahedron Lett. 2000, 41, 7851–7853.

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214. Ludvík, J.; Urban, J.; Jirkovský, J.; Zuman, P. Evidence of N-N single bond as a hindrance of electron delocalization in cyclic and acyclic azines. In Reactive Intermediates in Organic and Biological Electrochemistry; Yoshida, J.; Peters, D. G.; Workentin, M. S. eds., The Electrochemical Society—PV, 2001, Issue 14, Pennington, NJ, pp. 132–135. 215. Ludvík, J.; unpublished results, to be submitted. 216. Glaser, R.; Chen, G. S. J. Comput. Chem. 1998, 19, 1130–1140. 217. Lewis, M.; Barnes, C. L.; Glaser, R. Can. J. Chem. 1998, 76, 1371–1378. 218. Saied, T.; Benkhoud, M. L.; Boujlel, K. Synth. Commun. 2002, 32, 225–233.

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32

Reductions of Carboxylic Acids and Derivatives Rolf Breinbauer and Martin Peters

CONTENTS I. Introduction ......................................................................................................................... 1249 II. Cathodic Reduction of Carboxylic Acids ............................................................................ 1249 III. Cathodic Reduction of Esters and Peresters........................................................................ 1251 IV. Cathodic Reduction of Anhydrides and Acyl Halides ........................................................ 1255 V. Cathodic Reduction of Amides, Lactams, Imides, Imidates, and Hydrazides ................... 1256 VI. Cathodic Reduction of Nitriles and Acyl Cyanides ............................................................ 1260 VII. Cathodic Reduction of Thioesters and Thio-Derivatives .................................................... 1261 References .................................................................................................................................... 1262

I. INTRODUCTION This chapter is focused on the cathodic reduction of carboxylic acids and its derivatives. Included are carboxylic acids, esters, peresters, lactones, and anhydrides. Because of covering the compound class of acyl halides in Chapter 25, the acyl halides are only briefly discussed, to put them into the context of other acyl derivatives. Also included are the cathodic reduction of amides, lactams, imides, imidates, and hydrazides. Section VI is focused on the reduction of nitriles, but also isocyanates, cyanohydrins, and cyanides are included. The chapter closes with a short discussion of sulfur analogs of carboxylic acids, especially thioesters and their derivatives. Good reviews and compilation of electrochemical research in the last century have been published recently [1–3]. Cathodic transformation of carboxylic acids and their derivatives are well known [4,5], and in the last decade the mechanistic understanding about the electrode processes has been tremendously improved, which has stimulated more applications [4–13].

II. CATHODIC REDUCTION OF CARbOXyLIC ACIDS A considerable challenge in the reduction of carboxylic acids is to stop the reduction at the aldehyde oxidation level as these compounds can be easily reduced further to alcohols or even alkanes (Scheme 32.1) [14–17]. The reduction of aliphatic acids and unactivated aromatic acids results in the production of hydrogen gas and the carboxylate anion [4,18]. In aqueous solution, the aldehyde reaction product predominantly exists as the hydrate, which is reduced at lower potentials than the carboxylic acid starting materials. For example, cinnamic acid can be reduced to cinnamaldehyde using either Zn or amalgamated copper [19]. In a similar manner, the reduction of oxalic acid leads to glyoxylic acid [20–22].

1249

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1250

Organic Electrochemistry O H

R

+2 e– +2 H+ –H2O O O

R

–

+ H2

+e–, +H+

+4e–, +4H+

O

R

–H2O

OH

R

OH

+6 e– +6 H+ –2 H2O R

CH3

SCHEME 32.1 Possible reactions in the cathodic reduction of carboxylic acids. (From Wagenknecht, et al., in: Lund, H., Hammerich, O., eds., Organic Electrochemistry; 4th edn., Marcel Dekker, Inc., New York, 2001, pp. 453–470.)

Wagenknecht et al. have defined the following requirements that must be fulfilled for the reduction of a carboxyl group to take place [4]: • The carboxyl group must be activated by an electron-withdrawing group. • The carboxyl group must be the most easily reducible group in the molecule. • The formed aldehyde must be protected against further reduction by formation of a nonreducible derivative, such as a hydrate or hemiacetal, borate complex, or bisulfite adduct [4]. Ohmori et al. have devised a sophisticated strategy in which triphenylphosphine (1) as an additive gets first oxidized at the anode, allowing the formation of acyloxytriphenylphosphonium ions (2), which are simultaneously reduced at the cathode to triphenylphosphine oxide (3) and aldehyde (Scheme 32.2) [23]. With keto acids and tributylphosphine (Bu3P) as substrates, the corresponding α-hydroxycycloalkanones are produced in good yield [24].

Anode

Cathode +

Ph3P

+

O –e– –H+

–e–

O

Ph3P

RCOOH

R

2

+e–

Ph3P 1

O Ph3P = O

+

R 3 +e– O H

O

+H+ R

–

R

SCHEME 32.2 Triphenylphosphine (1) additive allows the selective reduction of carboxylic acids to aldehydes. (From Maeda, H. et al., Tetrahedron Lett., 33, 1347, 1992.)

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1251

Reductions of Carboxylic Acids and Derivatives OH

CO2H

Pb cathode H2O/NH4OH/(NH4)2CO3 80%

CO2H

CO2H

4

5

SCHEME 32.3 Chemoselective reduction of terephthalic acid (4). (From Wagenknecht, J.H. et al., eds., Organic Electrochemistry, 4th edn., Marcel Dekker, Inc., New York, 2001, pp. 453–470; Degner, D., Top. Curr. Chem., 148, 1, 1988.)

Aromatic carboxylic acids can be reduced to the corresponding benzylic alcohols. For example, BASF has reported about the selective reduction of terephthalic acid (4) to 4-(hydroxylmethyl)benzoic acid (5) in excellent chemical yield and a current efficiency of 48% (Scheme 32.3) [22].

III. CATHODIC REDUCTION OF ESTERS AND PERESTERS Depending on the substrate (owing to the individual stability of potential fragments) and the electrolysis conditions, the reduction of esters can lead to several products (Scheme 32.4). It is generally accepted that the first electrode reaction mechanism involves an electron transfer, which produces a radical anion [25–37]. From a synthetic organic perspective, both reaction pathways in the cleavage of esters are of synthetic interest. Marko et al. have reported high yields for preparative reactions in which either toluate esters 6 are reductively cleaved to the corresponding alcohols 7 (in the presence of a protic co-solvent) [38] or to the corresponding alkanes 8 (in aprotic electrolyte) (Scheme 32.5) [39]. Both reactions have been tested in highly functionalized substrates and are of considerable synthetic potential, especially the formal deoxygenation of alcohols in the second pathway, which competes well with alternative chemical procedures such as the Barton–McCombie reaction. A spectacular example for the selectivity and scalability of electrochemical ester reduction is the key step in a new synthetic route toward the antibiotic ceftibuten starting from readily accessible penicillin G [40]. In the synthetic sequence, allyl acetate 9 is reductively cleaved at an Sn cathode in a divided cell forming cepham 10, in which the double bond of the allyl moiety has shifted to an exocyclic position (Scheme 32.6). The process has been performed in 3 kg batches in >95% isolated yield and a current efficiency of 4–6% [41]. The deoxygenation of allylic, benzylic, and vicinal diols has been reported to occur favorably by the reduction of oxalate esters, as these show favorable redox potentials (E° −1.6 to −1.7 V vs. SCE) (Scheme 32.7) [42]. O R

– O

+e–

– O

R’

+

R’

H

R’

O

R

O

R’

O

R

+H+ +e–

OH R

R’ O

+H+

OH R

O +

R’ O

R

H +2e– +2H+

R

SCHEME 32.4

General pathways in the reductive cleavage of esters.

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OH

HO

R’

1252

Organic Electrochemistry Graphite electrode divided cell 0.15M n-Bu4NBF4

R-OH

NMP[a]/iPrOH (93/7) 7

– O

O O

R O

R

6 Graphite electrode divided cell 0.15M n-Bu4NBF4 R-H NMP[a], 130ºC 8

SCHEME 32.5 Control over product selectivity by electrolyte composition ([a] NMP = N-methyl-2pyrrolidone). (From Lam, K. and Markó, I. E., Org. Lett., 11, 2752, 2009; Lam, K. and Markó, I. E., Chem. Commun., 95, 2009.) O S

H N

HO2C O

N

OAc

O CO2H

Sn mesh cathode 0.2M K2HPO4

HO2C O

N O CO2H

99% conversion >95% yield

9

O S

H N

10

SCHEME 32.6 Pilot scale ester reduction in the synthesis of the antibiotic ceftibuten. (From Bernasconi et al., Org. Process Res. Dev., 6, 158, 2002.)

O EtO

Ph O

O O

Hg cathode 0.1M n-Bu4NCIO4-DMF, 2 F

O

Ph 11

OEt O

80%

Ph H

H Ph 12

SCHEME 32.7 Stilbene (12) formation by reducing oxalate ester 11. (From Utley, J. H. P. and Ramesh, S., ARKIVOC (Gainesville, FL, U. S.), 18, 2003.)

Paired electrochemical syntheses, in which both the anodic and cathodic reactions can be matched and contribute to the formation of valuable products, represent a formidable goal in green chemistry [43]. In 1999, BASF introduced the first paired electrosynthesis in a technical application (Scheme 32.8) [44–46]. Methyl phthalate (13) is cathodically reduced to phthalide (14), which serves as intermediate for the synthesis of fungicides. Concomitantly, 1-(tert-butyl)-4-methylbenzene (15) is oxidized to give 1-(tert-butyl)-4-(dimethoxymethyl)benzene (16), which is another important fine chemical intermediate. This paired electrosynthesis not only shows current yields of up to 190% but also features perfect atom economy since in the undivided cell the MeOH produced in the phthalate reduction can be used for the formation of the dimethylacetal 16. The process is reported to afford 4000 t/a [47,48].

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Reductions of Carboxylic Acids and Derivatives Cathode

Anode CO2CH3

4e– CO2CH3 13

+4H+

15

+2MeOH

O H3CO O

4e–

OCH3 14

+2MeOH

16

+4H+

SCHEME 32.8 BASF paired electrosynthesis process for the production of phthalide (14) and dimethylacetal 16. (From Frontana-Uribe, B.A. et  al., Green Chem., 12, 2099, 2010; Frontana-Uribe, B. A., Proc. Electrochem. Soc., 2002–2010, 52, 2002.)

Ni-cathode divided cell

O N HN

OMe NH2 17

1M LiCl/EtNH2 6F 44%

N

OH NH2

HN 18

SCHEME 32.9 Reduction to l-histidinol (18) using electrochemically generated solvated electrons. (From Beltrá, A.P. et al., J. Electrochem. Soc., 152, D65, 2005.)

In contrast to the well-established reduction of aromatic esters, the electrochemical reduction of aliphatic esters represents a considerable challenge due to the highly negative reduction potentials of these substrates (−3.0 V vs. SCE). Several reports use electrolyte combinations in which solvated electrons are formed, capable of reducing aliphatic esters. For example, l-histidinol (18) can be prepared from l-histidine methyl ester (17) in an LiCl/EtNH2 electrolyte (Scheme 32.9) [49]. Through the use of Mg electrodes, aliphatic esters can be reduced to the corresponding alcohols in very good yields in THF/tBuOH as an electrolyte (Scheme 32.10) [31]. In the absence of a protic co-solvent, 1,2-diketones 19 are formed, whereas the addition of chlorotrimethylsilane (Me3SiCl) produces bis(trimethylsilyloxy)alkenes 20 [29]. In an interesting variant of this principle, the electroreduction of aliphatic esters 21a-b in the presence of diene 22 or styrene 23 produces the corresponding carbocycles 24a-b (Scheme 32.11) [36]. In a similar manner, olefinic esters can be used for an electroreductive intramolecular cyclization [50], which has been exploited in an interesting synthesis of muscone [51]. In a paired electrosynthesis of reduction of aliphatic esters at an Mg cathode and anodic oxidation of THF at a Pt anode, the corresponding tetrahydrofuranyl-protected alkanols are produced [50,52]. All these reactions have in common that they work only with Mg electrodes, suggesting that either Mg(0) generated by the cathodic reaction reacts as an active reductant creating a mediatory system or that Mg ions serve as Lewis acids that shift the reduction potential to more favorable terms. Kise et al. reported about an unusual head-to-tail coupling of alkylbenzoates 25 by electroreduction in an undivided cell with an Sn cathode in iPrOH containing tetraalkylammonium salt as a supporting electrolyte. Under the protic conditions, one of the benzene rings gets reduced to a cyclohexane derivative 26 (Scheme 32.12) [53].

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Organic Electrochemistry +e–

R

0.5M LiClO4 THF/tBuOH

O

+e– Mg electrodes

OMe

R

O

–

R

O

+e–

R

R

0.5M LiClO4 THF

OMe

OH

O 19

+e– R

0.5M LiClO4 THF/Me3SiCl

OSiMe3 R OSiMe3 20

SCHEME 32.10 Reduction to aliphatic esters in the presence of Mg electrodes. (From Kashimura, S. et al., Tetrahedron Lett., 36, 4805, 1995; Shono, T. et al., J. Org. Chem., 57, 1061, 1992.) O

+e–

+

Mg electrodes 0.5M LiClO4/THF

OMe 21a

OH

22

24a 71%

O

OH

+e–

+

Mg electrodes 0.5M LiClO4/THF

OMe 21b

82%

23

24b

SCHEME 32.11 Formation of carbocycles 24a–b by electroreductive coupling in the presence of Mg electrodes. (From Shono, T. et al., J. Org. Chem., 57, 5561, 1992.) OH

O

+e– OCH3

2

Sn cathode Et4NOTs/iPrOH

25

CO2CH3 26

73% O Via:

CO2CH3

SCHEME 32.12 Reductive head-to-tail coupling of alkyl benzoates 25. (From Kise, N. et al., J. Org. Chem., 66, 862, 2001.)

The issue of ester reduction versus arene nucleus reduction is especially relevant in the reduction of pyridine esters. By careful control of the electrolysis conditions, the product formation can be optimized to the desired product [54]. For example, one electron reduction of pyridine carboxylic alkyl esters in acetonitrile (ACN) leads to relatively stable radical anions that decompose via cleavage of the oxygen alkyl bond to form the carboxylate anion 28 and the alkyl radical (Scheme 32.13) [27]. The alkyl radical rapidly reacts to form hydrocarbons like propane or hexane (by dimerization).

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Reductions of Carboxylic Acids and Derivatives

N

– O

+e–

O

N

glassy carbon 0.2M Et4NBF4, ACN

O

O

27

28 93%

SCHEME 32.13 Cleavage of O-alkyl bond in reduction of pyridine carboxylic alkyl ester 27. (From Webster, R.D. and Bond, A.M., J. Org. Chem., 62, 1779, 1997.) Hg cathode H2O/H3BO3/K2SO4

HO O

O

HO O

87%

OH OH

OH

OH OH 30

29

SCHEME 32.14 Electrochemical production of d-ribose (30). (From Degner, D., Top. Curr. Chem., 148, 1, 1988.) O R

O

+e– O O

tBu

R

– O

+ tBu

O

32 31 tBu

– O

SCHEME 32.15 Reductive cleavage of peroxyester 31. (From Antonello, S. et al., J. Am. Chem. Soc., 123, 9577, 2001; Antonello, S. and Maran, F., J. Am. Chem. Soc., 121, 9668, 1999.)

Under different conditions, the pyridine nucleus can be reduced without affecting the ester moieties [55]. A mechanistic model has been established that accounts for all sequential reaction products generated in the reduction of ethyl picolinate [56]. The reduction of lactones occurs similarly to acyclic esters. A notable example is the cathodic reduction of d-ribono-γ-lactone (29) to d-ribose (30) (Scheme 32.14) [22]. The reduction of peresters results first in the cleavage of the O–O bond (see Chapter 14) and is followed by subsequent reduction of the alkoxy radicals (Scheme 32.15) [57,58].

IV. CATHODIC REDUCTION OF ANHyDRIDES AND ACyL HALIDES Phthalic acid anhydrides are reduced under aprotic conditions upon uptake of one electron in the range of −0.7 to −1.2 V (vs. SCE) depending on the substituents on the benzene ring. Cathodic reduction of phthalic acid anhydrides at an Hg cathode in the presence of excess Me3SiCl produced the corresponding diphthalic acid lactonether 34 in very good yields (Scheme 32.16) [59]. O O 2

O O 33

Hg cathode Et4NCl, ACN Me3SiCl 81%

TMSO

O O

OTMS

O 34

SCHEME 32.16 Reductive dimerization of phthalic acid anhydride (33). (Troll, T. and Ollmann, G.W., Tetrahedron Lett., 22, 3497, 1981.)

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Organic Electrochemistry

+e– O R

O–

+e– Cl

R

R

O – +H+

O

O R

R

Cl

35

H

36

O R

R O 19

SCHEME 32.17

Possible reaction pathways in the reduction of acyl chlorides 35.

Cathodic reduction of acyl chlorides typically results in cleavage of the C–Cl bond (see Chapter  25) to generate an acyl radical 36 and a halogenide anion. Depending on the reaction conditions, the acyl radical can be transformed to an aldehyde by abstraction of a proton from the solvent or further reduced to the corresponding anion, which accepts a proton from residual water or supporting electrolyte. Sometimes also 1,2-diketones 19 (or derivatives thereof) resulting from the dimerization of the acyl radical are observed as the main product (Scheme 32.17) [60–66]. In the presence of a Ni-salen electrocatalyst, the reduction of cyclohexanecarbonylchloride does not produce any aldehyde, but instead the tetrameric product 1,2-dicylohexene-1,2-diol dicyclohexanoate in excellent 83% yield [67].

V. CATHODIC REDUCTION OF AMIDES, LACTAMS, IMIDES, IMIDATES, AND HyDRAZIDES The cathodic transformation of amides, lactams, and imides is normally performed at mercury or lead electrodes. Under strongly acidic conditions, the reduction of the carbonyl group to the corresponding methylene group is usually observed. However, under mild acidic conditions also cathodic cleavage may occur [68,69]. Voltammetric and EPR studies have shown that amides can be reversibly reduced to relatively stable anion radicals [12,70–72]. The reduction of amides, lactams, imides, imidates, and hydrazides occurs through two-electron reduction processes. The reduction of an amide can either proceed via an aminal intermediate to a tertiary amine, or lead via reductive C–N cleavage to the corresponding aldehyde, which can be further reduced to deliver an alcohol and a secondary amine (Scheme 32.18) [73,74]. The carbonyl-function of the amide moiety is reduced under strongly acidic conditions with a Pb or Hg cathode under formation of the methylene function. Electrochemical reduction of aliphatic amides with an Mg cathode in the presence of a proton-donor as a co-solvent (tBuOH) leads to the corresponding aliphatic alcohols. If a smaller amount of proton-donor (3.5 eq tBuOH) is used, the reduction of aliphatic amides can be controlled to stop at the aldehyde stage [31]. The reduction of amides 37 in the presence of proton-donor (tBuOH) and Me3SiCl (undivided electrolysis cell, Mg electrodes, N2 atmosphere) gave α-amino-ketones 38 in good to very good yields, after hydrolysis of the resulting enolether 39 (Scheme 32.19) [75]. 3-Nitrobenzoyl groups can be used as protecting groups for aliphatic amines. The nitrobenzamides could be cleaved at the C–N bond in DMF/0.1M n-Bu4NBF4 solution [76]. Even in the

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Reductions of Carboxylic Acids and Derivatives –

O R2

R1

N R3

+e– OH R1

O

+2e–, +2H+

R2

R2

R1

N R3

+2e–, +2H+

N R3

R2

O R1

HN R3

+2e–, +2H+ R1

OH

+H+, –H2O R2

+

R1

+2e–, +2H+

N R3

H R2 N+ R3

R1

SCHEME 32.18 Possible reaction pathways of reducing amides and its derivatives. (From Wagenknecht, J.H. et al., in: Lund, H., Hammerich, O., eds., Organic Electrochemistry, 4th edn., Marcel Dekker, Inc., New York, 2001, pp. 453–470.) O R1

Me3SiCl, LiClO4, THF Mg electrodes

NR2

OTMS

O R1

R1

10% HCl

R1

R1

NR2

NR2 39

37

38

SCHEME 32.19 Reduction of amides in the presence of Me3SiCl. (From Kashimura, S. et al., Tetrahedron Lett., 39, 6199, 1998.)

presence of a peptidic macrolactone function (antibiotic pristinamycin IA) a selective cleavage of the R(CO)–NHR′ bond is feasible at an Hg cathode under aqueous acidic conditions, obtaining the corresponding alcohol RCH2–OH and the free amine R′–NH2 [77–81]. While lactams can be reduced by metal-hydrides in a nonelectrochemical reaction, the use of lactams in electrochemistry is more centered on its application as a probase (Scheme 32.20, see Chapter 43). After electrochemical reduction, the corresponding electrogenerated base (EGB) can be used as a strong Brønsted base [82–86]. An interesting application is the N-acryloylation of

NH probase

O

+e–, ACN, Et4NClO4

O

divided cell, Pt electrodes, N2 atmosphere, galvanostatic control, I = 25 mA/cm2, 1 F

N – electrogenerated base (EGB)

40

41 O +

O HN

O

O +

R4 R3

R1

R2 42

X X

R4

N – R3

O

O N

R1

O R2

X = Cl, Br

SCHEME 32.20 Application of 2-pyrrolidone (40) as probase in the N-acryloylation of chiral auxiliaries 42. (From Feroci, M. et al., Eur. J. Org. Chem., 2765, 2001.)

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Organic Electrochemistry

chiral oxazolidin-2-ones. As probase 2-pyrrolidone (40) is electrochemically reduced under N2 atmosphere to the corresponding 2-pyrrolidone anion (41) (ACN, 0.1M Et4NClO4, Pt electrodes, galvanostatic, I = 25 mA/cm2) and is used in the deprotonation of chiral Evans’ auxiliaries 42 (oxazolidin-2-ones) [87]. Phthalimides can be easily reduced and many investigations have been performed for mechanistic studies [88–93]. In aprotic or alkaline media, two one-electron reduction steps are observed in voltammetric experiments [94–96]. By double decarbonylation of phthalimide 43 in a water– ACN mixture, the isoindoline (44) could be formed; in alkaline solution, phthalimide is reduced to dimeric epoxide 45 (Scheme 32.21) [95,97,98]. EPR-analyses of N-substituted phthalimides showed two one-electron processes. The initially formed dianion is protonated by a proton-donor (like phenol or protic solvent). It is reasonable that the tautomer 3-hydroxyphthalimidine is in equilibrium with the open aldehyde 47. After reduction of aldehyde 47, either the corresponding alcohol 48 or the phthalide 14 is formed by cyclization (Scheme 32.22). A synthetic approach of synthesizing silyl ketene acetals is reported by Kise et al. After electroreduction, an intramolecular coupling of the α,β-unsaturated moiety of phthalimides 49 (in the

O

HN

O

NH

O

base +H2O, +2e–

NH

–2 OH– –1.2V vs. SCE

O

O 45

43

+4e–, +4H+

NH

–H2O ACN/H2O = 2/1 pH = 2 –1.1 V vs. SCE

O +4e– +4H+ –H2O

NH 44

SCHEME 32.21 Reduction of phthalimide (43) under acidic and basic conditions. (From Wagenknecht, et  al., in: Lund, H., Hammerich, O., eds., Organic Electrochemistry; 4th edn., Marcel Dekker, Inc., New York, 2001, pp. 453–470; Leedy, D. W. and Muck, D. L., J. Am. Chem. Soc., 93, 4264, 1971; Porter et al., J. Electrochem. Soc., 126, 1693, 1979; Fechete, I. and Jouikov, V., Electrochim. Acta, 53, 7107, 2008.) OH +H+ H N – O

O

O +2e–, +2H+

H N

N Ph O 46

O

Ph

O

+2e–, +H+

48 H N

Ph

Ph

O

47

O –aniline

O 14

SCHEME 32.22 Reduction of N-substituted phthalimide 46. (From Wagenknecht, J.H. et al., in: Lund, H., Hammerich, O., eds., Organic Electrochemistry, 4th edn., Marcel Dekker, Inc., New York, 2001, pp. 453–470; Farnia, G. et  al., J. Electroanal. Chem., 33, 31, 1971; Leedy, D.W. and Muck, D.L., J. Am. Chem. Soc., 93, 4264, 1971; Ryvolova-Kejharova, A. and Zuman, P., Collect. Czech. Chem. Commun., 36, 1019, 1971.)

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1259

Reductions of Carboxylic Acids and Derivatives OTMS COOEt

O

Divided cell, Pt electrodes

n

N

H TMSO

n

0.3M Et4NOTs, DMF, Me3SiCl, NEt3, 3.1 F

O

OEt N

n = 1, 2 49

O 50

TBAF, THF, 0°C

COOEt H TMSO n N O 51

SCHEME 32.23 Stereospecific intramolecular trans-cyclization of phthalimides 49 in the presence of Me3SiCl. (From Kise, N. et al., Org. Lett., 11, 4902, 2009.)

presence of Me3SiCl and Et4NOTs) results in silyl ketene acetals 50, which can be desilylated by treatment with TBAF in THF; after treatment stereospecific trans-cyclized products 51 are observed (Scheme 32.23) [99]. N-acylated imidates can be electrochemically dimerized at an Hg electrode in aprotic solvents [100,101], dicarboximides (like N-substituted oxazolidine-2,4-diones) can irreversibly be reduced at −2.2 V vs. Ag/AgCl [102]. Most commonly used in the reduction of hydrazides is the hydrogenolysis with Raney nickel or Pt. The reduction of carboxylic acid hydrazides to amides in good up to very good yields has been described [103] using an Sn cathode, graphite anode, and Britton–Robinson buffer (25% EtOH, pH 5.5) in a divided cell [104–111]. The reduction of carboxylic acid hydrazides 52 is feasible (potential difference 2.5 V), even in the presence of aryl halogen or olefinic groups (Scheme 32.24) [112]. Trace measurements of isoniazide in human urine with screen-printed carbon electrodes modified with poly-l-histidine [113] or by isoniazide determination in medicines by differential pulse voltammetry, capillary electrophoresis, or polarographic reduction methods are known [114–119].

O R

N H

NH2

O

+2e–, +2H+ – NH3

R

NH2

R = phenyl, alkyl, cyclohexyl 52 O

N

N H

NH2

O

+2e–, +2H+ – NH3

+4e–, +4H+ NH2 N

– NH3, – H2O

OH N

SCHEME 32.24 Reduction of carboxylic acid hydrazides 52. (From Mentel, M. et al., Synthesis, 9, 1463, 2009.)

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VI.

Organic Electrochemistry

CATHODIC REDUCTION OF NITRILES AND ACyL CyANIDES

Raney nickel can be used as a catalytic electrode for the electrohydrogenation of nitriles to synthesize the corresponding primary amines [120–122]. It is remarkable that the electrochemical reduction is possible at room temperature at atmospheric pressure, whereas the conventional Raney nickel hydrogenation usually requires high pressure and temperature [123]. Decyanation under electrochemical conditions in anhydrous amine media was first reported by Arapakos and Scott [124]. Under acidic conditions, the reduction of nitriles resulted in the corresponding amines, whereas reducing nitriles under neutral or basic conditions, decyanation is preferred, by cleavage of the C–CN bond to form the corresponding hydrocarbon derivatives (Scheme 32.25) [4,125–127]. The reductive decyanation under electrochemical conditions was used in the total synthesis of (±)-hirsutene (53) (LiCl, EtNH2, and Pt electrodes) (Scheme 32.26) [128]. The adsorption phenomena and the mechanism of electrochemical reduction of cyano-pyridines or aromatic nitriles have been the subject of intensive investigations [129–132]. One interesting transformation of nitriles is depicted in Scheme 32.27. 4-Cyanocinnolines 54 can be transformed to 4(1H)-cinnolones 55 in the presence of O2 and water [133]. Acetonitrile can be reduced to the cyanomethyl anion in the absence of a proton donor. The Et4NClO4 stabilized cyanomethyl anion attacks a parent molecule ACN producing the 3-aminocrotonitrile anion. The electrogenerated anions can react in the presence of CO2 with α-haloamides to the corresponding oxazolidine-2,4-diones [134]. In the presence or absence of water, benzoyl cyanides 56 can be reduced to 1,2-diketones 19 or mandelonitrile benzoates 57 (Scheme 32.28) [123,135,136]. Nevertheless, the resulting mandelonitrile benzoates 57 can be reduced in an aprotic media to the corresponding vicinal diketones 19 (Scheme 32.29) [137].

NH2

+4e–, +4H+

R

SCHEME 32.25

R CN

+e– R CN

–

R

+

CN–

+e–, +H+

R H

Reduction of nitriles under acidic or neutral/basic conditions.

HO

HO Decyanation

NC

H

H

H

LiCl, EtNH2, Pt

2 steps

H

H H

H

H

53

SCHEME 32.26 Electrochemical decyanation in the synthesis of natural product (±)-hirsutene (53). (From Franck-Neumann, M. et al., Tetrahedron, 48, 1911, 1992.)

O

CN R N N 54

R

+2e– Pt electrodes, PhCN, n-Bu4NBF4, O2, H2O

N N– H 55

SCHEME 32.27 Reduction of 4-cyanocinnolines 54 to 4(1H)-cinnolones 55. (From Matsubara, Y. et al., Tetrahedron Lett., 41, 7901, 2000.)

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Reductions of Carboxylic Acids and Derivatives

CN Ar

O

O

O

Et4NOTs/ACN, 0.4 A, 1.2 F

Ar

Ar

57

O

Et4NOTs/H2O/ACN,

2

Ar

Ar

0.4 A, 1.2 F

CN

O 19

56

SCHEME 32.28 Reduction of benzoyl cyanides 56. (From Okimoto, M. et al., J. Org. Chem., 61, 4835, 1996.)

CN Ar

O

O

CN

base

O

Ar – O

Ar

O

NC Ar Ar

–

O

–CN–

Ar

Ar

Ar

O

O 19

57

SCHEME 32.29 Reduction of mandelonitrile benzoates 57. (From Zheng, Z.-R. and Lund, H., J. Electroanal. Chem., 441, 221, 1998.)

VII. CATHODIC REDUCTION OF THIOESTERS AND THIO-DERIVATIVES Aromatic and aliphatic thioesters 58 can be regioselectivly cleaved depending on the nature of the substituents. The carbonyl–sulfur cleavage leads to the corresponding acyl radical 59, and cleavage between the RCOS group and the alkyl group afforded thiocarboxylic acid anion 60 (Scheme 32.30) [138]. However, aliphatic thioesters are difficult to reduce due to their low reduction potentials (−3.0 V vs. SCE) [4]. Thioesters like RCOSAr (R = alkyl) undergo chemoselective cleavage of the carbonyl–sulfur bond; in the presence of DMF, the corresponding RCONMe2 are synthesized. Thioamidoimidates 61 can be reduced in aprotic solvent at a mercury pool electrode to the corresponding thiazolo[5,4-d]thiazoles 62 by dimerization (Scheme 32.31) [139]. Cyclic ketones 64 can be synthesized by reduction of S-(2-methoxycarbonyl)phenyl thioesters 63 (Scheme 32.32) [140]. The voltammetric behavior of thioic S-esters and dithioic S,S′-diesters has recently been studied. In acetonitrile, the functional group is cleaved at the C(O)–S bond as a one-electron reduction

O R1

S

+e–

R2

O R1

– S

58

O

carbon fiber circular cathode, Mg anode, R2

LiCIO4, DMF, N2 atmosphere, 60 mA, 0.3 A/dm2, 2.2 F

–

+

R1

S

R2

59 O–

O – R2

R1

S

–

R1

S

60

SCHEME 32.30 Two possible pathways for the reduction of thioesters 58. (From Weïwer, M. et  al., Tetrahedron, 61, 1709, 2005.)

R1 R3

O

S N

mercury pool electrode, N H

R2

DMF, n-Bu4NBF4, N2 atmosphere R1, R2, R3 = Ph, benzyl, Me

61

HN

R2

N

S

S

N

R1

R1 R2

NH 62

SCHEME 32.31 Reduction of thioamidoimidates 61 produces thiazoles 62 by dimerization. (From Tapsoba, I. et al., J. Electroanal. Chem., 569, 89, 2004.)

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1262

Organic Electrochemistry graphite plate cathode, aluminum rod anode,

O R2

X

n S

R1

63

COOMe

DMF, Et4NCIO4, N2 atmosphere, 6 mA, 4 F

O R2 X

1 n R 64

O +

R2

X R1

n

1, R2 = H, Ph, Me

R

X = CH2, O n = 1, 2

SCHEME 32.32 Syntheses of cyclic ketones 64 by electroreduction of S-(2-methoxycarbonyl)phenyl thioesters 63. (From Ozaki, S. et al., J. Org. Chem., 66, 2503, 2001.)

step between −1.61 to −2.69 V vs. Fc/Fc+ reduction potential [27,28,33,35,141]. The reduction of thionicotinamide and thioisonicotinamide in strongly acidic media (pH < 4) proceeds through an amino-thiol intermediate after two one-electron processes [142]. Benzanellated sulfur heterocycles can be synthesized by electroreduction of dithiocarboxylic esters with leaving groups at the benzene ring [143].

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33

Oxidation of Carboxylic Acids and Derivatives Hideo Tanaka, Manabu Kuroboshi, and Sigeru Torii

CONTENTS I. Introduction ........................................................................................................................... 1267 II. Experimental Conditions for Kolbe and Non-Kolbe Reactions ............................................ 1269 A. Constant-Current-Density Electrolysis (Electrode Potential and Current Density) ..... 1269 B. Electrode Material ........................................................................................................ 1270 C. Solvent, Electrolyte, and Additives ............................................................................... 1270 D. Other Variables (Temperature, Pressure, and Magnetic Field)..................................... 1271 E. Conclusion ..................................................................................................................... 1272 III. Radical Reaction (One-Electron Oxidation) ......................................................................... 1272 A. Kolbe Coupling Reaction .............................................................................................. 1272 B. Kolbe Cross Coupling Reaction .................................................................................... 1276 C. Decarboxylative Radical Addition ................................................................................ 1279 D. Miscellaneous ............................................................................................................... 1283 IV. Carbenium Ion Reaction (Two-Electron Oxidation)............................................................. 1285 A. Acyloxylation, Alkoxylation, Hydroxylation, and Acetamidation................................1286 B. Alkene Formation..........................................................................................................1291 C. Rearrangement ..............................................................................................................1296 D. Formation of Aldehydes, Ketones, Acetals, and α,β-Enones ........................................1299 E. Miscellaneous ............................................................................................................... 1301 References .................................................................................................................................... 1302

I. INTRODUCTION The coupling reaction of alkyl radicals provided by electrodecarboxylation of alkanoates R–CO2− gives the Kolbe dimer R–R (Equation 33.1). 2RCO2–

–2CO2 –2e–

R–R

(33.1)

Though the Kolbe reaction sounds historical in electroorganic chemistry, it is still alive as an important synthetic tool in today’s organic synthesis. Since the discovery of the Kolbe reaction [1], a vast number of investigations have been carried out to unveil the features of its synthetic potentiality. The first finding of the versatility by Wurtz [2] is that the electrolysis of two different carboxylates R1–CO2− and R2–CO2− gives the mixed dimer R1–R2 together with two possible symmetrical dimers R1–R1 and R2–R2 (Equation 33.2). R1–CO2– + R2–CO2–

–2CO2 –2e–

R1–R2 + R1–R1 + R2–R2

(33.2)

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The most versatile extension of the Kolbe reaction was done by Brown and Walker [3], which opened a simple entry to the synthesis of a variety of dicarboxylic acids (Equation 33.3). Actually, this modification, the Brown–Walker reaction, is a well-established method for the synthesis of long-chain carboxylic acids (Equations 33.4 and 33.5). 2 RO2C–(CH2)n–CO2–

–2CO2

RO2C–(CH2)2n–CO2R

–2e–

–2CO2

R1–CO2– + R2O2C–(CH2)n–CO2–

–2e–

R1O2C–(CH2)m–CO2– + R2O2C–(CH2)n–CO2–

R1–(CH2)n–CO2R2

–2CO2

R1O2C–(CH2)m+n–CO2R2

–2e–

(33.3) (33.4) (33.5)

Brown and Walker first proposed the generally accepted mechanism of the Kolbe reaction, which involves the initial discharge of carboxylates at the anode followed by decarboxylation and subsequent combination of the resulting radicals, leading to the Kolbe dimer [3]. The radical formed also undergoes disproportionation to afford alkenes and alkanes (R–H) as the result of hydrogen abstraction (Scheme 33.1). In the 1960s, a more promising expansion of the electrodecarboxylation reaction in terms of synthetic utility has been recorded in which a carbenium intermediate R+ is formed by further one-electron oxidation of radical R• at the anode and plays an important role [4]. Depending on the structural characteristics of the carboxylates and/or electrochemical variables, the cation intermediate R+ undergoes the so-called non-Kolbe reactions, for example, substitution, deprotonation, C–C bond cleavage, and rearrangement to provide alcohols, ethers, esters, amides, alkenes, and others (Scheme 33.1). The electrochemical hydroxylation and esterification in aqueous media, the so-called Hofer–Moest reaction, initially are regarded as a side reaction, but sometimes become the major reaction in the aqueous Kolbe reaction [5]. Such puzzling phenomena encountered in the early Kolbe reaction are rationalized on the basis of the assumption just described earlier. In the past five decades, more sophisticated electrolysis conditions suited for the Kolbe dimerization, for hydrogen abstraction, for substitution, for rearrangement, and for other radical- or cation-induced reactions have been accumulated for a number of carboxylic acids and used for many synthetic purposes. Most recently, in the field of the two-electron transfer process, significant progress has been made in exploiting preparative use of the cation intermediates for the functionalization, such as acyloxylation, alkoxylation, hydroxylation, acetamidation, formation of alkenes and carbonyl groups, and rearrangements. The concept of modern electrodecarboxylation of carboxylic acids tends to be far from that of the old-fashioned Kolbe reaction. Nowadays, electrodecarboxylation is available as a potent tool for building up the functional groups in complex molecules. Indeed, recent progress in this field enables us to design various synthetic equivalents to, for example, carbonyl groups, α,β-enones, and double bonds by virtue of electrodecarboxylation. Today’s synthetic potentialities of the electrogenerated carbenium intermediates R+ are discussed in Section IV. R–CO2–

–e–

R–CO2

–CO2

–e–

R+

R +Nu–

Alkene, R–H

R–R

R–Nu

–H+ Alkene

Nu: OH, OMe, OAc, NHC(O)Me, etc.

SCHEME 33.1

One-electron oxidation versus two-electron oxidation.

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C–C bond cleavage rearrangement

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Oxidation of Carboxylic Acids and Derivatives

II.

EXPERIMENTAL CONDITIONS FOR KOLbE AND NON-KOLbE REACTIONS

Electrodecarboxylation has been examined in a variety of electrolysis systems, and a proper combination of solvent, supporting electrolyte, electrode material, pH, concentration of salt and substrate, and other variables has been chosen according to the synthetic requirements in each case. As briefly mentioned before, the electrodecarboxylation provides two competing pathways that give rise to either radical R• or cation R+ intermediates (Scheme 33.1), leading to various products. The optimum conditions for promoting a particular reaction have been accumulated. Thus far, one can, in many cases, design a suitable electrolysis system making the reaction products selectively via either the radical R• or cation R+ pathway.

A.

CONSTANT-CURRENT-DENSITY ELECTROLYSIS (ELECTRODE POTENTIAL AND CURRENT DENSITY)

The Kolbe reaction is usually performed with constant current density, and control of the anode potential by using a potentiostat is seldom essential. A high current density, higher than 0.25 and often up to 1.0 A/cm2, is in general recommended for the Kolbe dimerization, since the rate of the radical combination should be proportional to the concentration of the radical species per unit area of the electrode. It has been reported that if the anode potential is plotted against the logarithm of the current density (Tafel plots, Figure 33.1), there is a special value of current density at which the anode potential jumps from 2.1 to 2.4 V [6–8]. At a current density higher than that causing such a potential jump, discharge of the carboxylates proceeds almost exclusively and completely suppresses undesirable reactions that normally take place at a lower potential, such as oxygen evolution from water. The critical potentials at which the decarboxylation starts have been compiled for a variety of carboxylic acids; it is normally observed in a region of 2.0–2.8 V (vs. normal hydrogen electrode [NHE]) [9]. In practice, however, it is not necessary to control the anode potential, since the anode potential shifts to higher values than the critical potential even if the electrolysis is operated at a constant current density of around 1 mA/cm2. The radical R thus generated undergoes a further one-electron oxidation to give a cation R+. Owing to the high discharge potential of carboxylates, easily oxidizable radicals R• are believed to couple prior to the second electron transfer. Easily oxidizable functional groups such as amine, sulfide, and formyl, attached to the carbon atom bearing a carboxyl group, cannot survive under the electrolysis conditions.

Anode potential (V)

2.6 2.4 2.2 2.0 1.8 –1.0

0.0 1.0 log (current density/mA cm–2)

2.0

FIgURE 33.1 Anode potential (vs. NHE) as a function of log (current density) in an H2O–AcONa (0.5 M) system.

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1270

B.

Organic Electrochemistry

ELECTRODE MATERIAL

The influence of anode materials on the electrodecarboxylation in aqueous solutions has been well documented; platinum, iridium, and vitreous carbon electrodes favor the coupling product, whereas graphite and porous carbon anodes facilitate the two-electron process, giving cation intermediates. In contrast, in nonaqueous solvents, a variety of anode materials are employed. The choice of electrode material is less critical, although platinum is still recommended for the one-electron process and carbon for the two-electron process. Other electrode materials, such as gold, palladium, lead dioxide [10], iron-impregnated carbon [11], and shungite (a nongraphitized carbon) [12], are useful in some cases for the Kolbe coupling. The electrolysis of hydrophobic carboxylates such as long-chain or fluorinated ones is substrate-selectively suppressed on a PTFE composite-plated (hydrophobic) platinum anode (PTFE/Pt) [13]. The mechanism of the unique hydrophobic effect is discussed on the basis of preparative electrolysis, voltammetry, and hydrophilic/hydrophobic interaction between the anode surface and the carboxylate ions. The nature of the cathode material is not critical in Kolbe reactions. In most cases, neither starting material nor products are reducible at potentials less negative than that required for hydrogen evolution.

C.

SOLVENT, ELECTROLYTE, AND ADDITIVES

Neutral and slightly acidic media have been recommended for the Kolbe reaction. To keep the electrolyte acidic during electrolysis, the reaction is simply performed in an aqueous solution of the acid, neutralized by an alkali metal hydroxide or trialkylamine to an extent of 2–5%. Since hydroxyl ions are formed at the cathode at the same rate as carboxylate ions are consumed at the anode, the concentration of carboxylate ion remains approximately constant during the whole run. Kolbe coupling process of insoluble organic acids in an aqueous environment and with in situ emulsification by ultrasound is conducted cleanly and is highly charge efficient [14,15] (Equation 33.6). The electrode material used and the conditions employed during the electrolysis have only a slight effect on the type of products formed. Boron-doped diamond electrodes may be employed, replacing the less corrosive and ultrasound-resistant platinum electrodes. O 2R O–

Aq. NaOH (1 M), ultrasound (BDD)–(Pt), undiv. cell

R = nC5H11, nC6H13

R–R 15–45%

(33.6)

BDD: boron-doped diamond electrode

Water is now seldom used as a solvent in Kolbe oxidations. Methanol and aqueous methanol are regarded as the best choice as a solvent [16]. Dimethylformamide (DMF) and wet acetonitrile can also be used in particular cases [17–19]. Special attention should be paid to assess the solvent system for the generation and reaction of carbenium intermediates R+ under electrolysis conditions [16]. A variety of solvent systems elaborated for the electrodecarboxylation have been compiled [20]. On the basis of a rough statistical treatment of experimental findings, the combinations or interrelations among three kinds of solvent (water [W], protic solvent [P], and dipolar aprotic solvent [DA]), two reactive intermediates (cation R+ and radical R•), and four typical reactions (Kolbe dimerization, substitution, rearrangement, and alkene formation) are schematically illustrated in Figure 33.2 [20]. One can select the first choice of the solvent system that may be suited to undertaking certain reactions. For example, a dipolar aprotic solvent–water (DA–W type), such as pyridine–water, should be the first choice when aiming to obtain alkenes as the exclusive electrodecarboxylation product. Most of the electrodecarboxylations have been carried out with partially neutralized carboxylic acids. Alkaline and alkaline earth metal as well as ammonium (pyridinium) carboxylates work efficiently as supporting electrolytes. Recently, silica gel–supported piperidine also works as a base

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1271

Oxidation of Carboxylic Acids and Derivatives

W

Kolbe dimerization

Radical

Alkene formation DA –W

P–W

DA

P

Cation

Rearrangement

DA–P

Substitution

FIgURE 33.2 Correlation between solvent, reactive intermediate, and reaction type: W, water; P, protic solvent; DA, dipolar aprotic solvent. 2 R–CO2H

O

H N

CO2H

SiO2

N (0.1M)

MeOH/MeCN or MeOH (Pt)–(Pt), undiv. cell

R–R

O

44–99% H N

OMe

(Other 5 examples)

quant.

SCHEME 33.2

Kolbe and non-Kolbe reactions using silica gel–supported base.

for Kolbe and non-Kolbe reactions, wherein the solid-supported base is easily separated from the electrolytes and repeatedly used [21–23] (Scheme 33.2). Addition of certain salts, such as perchlorate, fluoroborate (BF4−), sulfate, dihydrogenphosphate (H2PO4−), hydrogencarbonate (HCO3−), and fluoride, tends to inhibit the radical reaction and favor the formation of cation intermediates [24–27]. The remarkable effects of the salts are well explained in terms of competitive adsorption between the anions and carboxylates. On addition of ammonium thiocyanate (NH4SCN), peracid R–C(O)OOH is obtained from R–CO2H [28]. The high oxygen-evolution overvoltage caused by NH4SCN makes potential domains for the discharge of water molecules and acetate ions overlap, leading to the oxidation of acetate ions, which involves water molecules and yields peracetates.

D. OTHER VARIABLES (TEMPERATURE, PRESSURE, AND MAGNETIC FIELD) Temperature is usually not a critical variable but improves viscosity and mass transport. An increase in temperature generally results in increase in yield of the Kolbe dimer [29,30]; however, over 50°C, in some cases, a total change in the product occurs [31,32]. Under high-pressure conditions, volatile species, such as hydrogen, carbon dioxide, and, in some cases, alkenes or alkanes, are accumulated on the anode surface, which may affect the product distribution [30,33]. An aqueous Kolbe electrolysis of an alkanoic acid (in the range of C4 –C6) tends to give higher yields of the Kolbe dimers when run at elevated pressure (100 kPa) than at atmospheric pressure. A magnetic field may affect the flow of electrons and ions, resulting in perturbation of the mass transport process near the electrode [34]. For example, in the decarboxylation of phenylacetic acid, the presence of a magnetic field increases the yield of aldehyde, presumably because the supply of molecular oxygen from the bulk solution of the electrode surface is enhanced.

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1272

Organic Electrochemistry

TAbLE 33.1 Effect of Experimental Variables on Electrodecarboxylation Variable

One-Electron Process

Two-Electron Process

Additive

High current density Pt electrode Acidic Neutral water (W) Protic (P) W−P None

Temperature

90%) retained.

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1275

Oxidation of Carboxylic Acids and Derivatives R3 R1

R3

–e– –CO2

CO2–

15–79%

R1 R2 11

R2 7 R2

R3 R1

R1

+

R1 R3

R2

(33.9) R3 R1 R2

R3 R1 R2 + R2

8

R1 R2 R3 10

R3 9

Kolbe oxidation of chiral hydroxyalkanoate gives the corresponding chiral diols in moderate yields [57]. The chiral center (β-position to the carboxylic acid bearing OH group) is retained during the electrolysis. Kolbe dimerization can be also applied to synthesis of disilanes (see Table 33.2) [60,61]. The selectivity depends on the structural parameters of the carboxylic acid and electrochemical conditions. The α-effect of silyl groups contributes to the stabilization of radicals to promote the formation of dimeric products. The electron donating nature of the silyl group causes the carboxyl group to be less acidic and more reluctant to form a carboxylate. Kolbe dimerization also gives fluorinated alkanes and (poly)ethers. The synthetic potentiality of the Kolbe dimerization has been well documented for a variety of symmetrical target molecules. For example, the dimerization of acids (12 and 13) is the key step of the pentacyclosqualene [70] and α-onocerin [71] syntheses. A large-scale production of sebacic acid has been realized by the Kolbe dimerization of the half-ester of methyl adipate [72]. CO2H

CO2H

O OAc HO 12

13

The Kolbe reaction of pyruvic acid and phenylglyoxylic acid (14) (R = Me and Ph) affords the corresponding diketones (15) in high yields (Equation 33.10) [73,74]. Interestingly, electroluminescence is observed during the electrolysis, which probably is attributed to phosphorescence due to the generation of excited triplets of α-diketones in the recombination of the acyl radical. RCOCO2– 14

MeCN–Et3N–(Pt)

RCOCOR + hν 15

–e– –CO2 [RCO ]

[RCO ]

R = Me (75%) R = Ph (90%)

(33.10)

[RCOCOR]*

Electrodecarboxylative polymerization of dicarboxylic acids has been examined [75–78]. Under the Kolbe electrolysis conditions, dicarboxylic acids (16) (n = 4–8), such as adipic acid and sebacic acid, give rise to polymer along with oligomeric carboxylic acids and oligomeric hydrocarbons (Equation 33.11). The polymer formation is explained by assuming an α,ω-biradical intermediate (17). HO2C–(CH2)n–CO2H

16

–2e– –CO2

(CH2)n

Polymer

(33.11)

17

The formation of grafted films on carbon electrode surfaces is attained by electrooxidation of arylacetates under Kolbe electrolysis conditions [79].

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1276

Organic Electrochemistry

B. KOLBE CROSS COUPLING REACTION Mixed Kolbe electrolysis of two different carboxylic acids gives a statistical mixture of the three coupling products. Thus, the cross coupling product is always accompanied by two symmetrical dimers owing to the statistical coupling of each radical intermediate R1• and R2• (Equation 33.12). Although a few attempts toward selective cross coupling—for example, square-pulse electrolysis of unsymmetrical acyl peroxide [80–82]—have been made, these are still far from the goal. The disadvantageous but unavoidable formation of the two symmetrical dimers can be reduced to a single product when one of them is used in 5–10-fold excess [83]. R1–CO2– + R2–CO2–

–2CO2

(33.12)

R1–R2 + R1–R1 + R2–R2

–2e–

The Kolbe electrolysis approach, even if not high yielding, presents some advantages over classical chemical methods, since it furnishes the desired compound in one single step and many different alkanoic acids can be used for the coupling reaction. Moreover, the reaction conditions are very mild so that many functional groups, for example, halogen, ketone, alcohol, ether, ester, lactone, and vinyl groups, are preserved during the electrolysis. Indeed, the Kolbe cross coupling may permit a variety of applications in natural product synthesis. By using this method, carotenoids, fatty acids, amino acids, muscone, fumulene [39], and pheromones [40] have been synthesized. Representative results of Kolbe cross couplings are compiled in Table 33.3 [83–111]. Enantiomerically pure chiral building blocks (18) for the synthesis of natural products are prepared by cross Kolbe electrolysis of chiral carboxylic acids 19 (Equation 33.13). Since radical coupling is very fast, that means with a very low activation energy, according to the Hammond principle, a very early transition state is to be expected, in which steric interaction between radicals 20 and 21 derived from electrodecarboxylation of 19 and R2–CO2H is still minimal. Nevertheless, the use of proven auxiliaries and large coradicals leads to a remarkable diastereoselectivity even in intermolecular radical coupling [112–114]. Representative results of Kolbe cross coupling of chiral carboxylic acids are compiled in Table 33.4 [104,114–122]. O Aux

CO2H + R2–CO2H

R1 19 Aux: chiral auxiliary

–e– –CO2

O + R2

Aux 20

R1

21

O Aux

(33.13)

O R2 R1

or Aux

R2 R1

18

Under the Kolbe cross coupling conditions, the C═C double bond of unsaturated carboxylic acids generally retains its configuration. Partial Z/E isomerization of alkenes is, however, observed in the electrolysis of (Z)-4-enoic acids (Equation 33.14) [123]. The electrolysis of (Z)-3-methyl-4octenoic acid (22) with acetic acid gives a mixture of Z and E isomers of the cross coupling product (23) (9:1–7:3) together with 24 and 25 (Equation 33.14). This can be well understood by assuming a Z/E isomerization via a reversible ring closure to cyclopropylcarbinyl radicals (26). CO2H

+

MeOH–KOH–(Pt)

MeCO2H

22 Me 23 (19–50%) E/Z = 9/1–7/3

+

24 (5–30%)

26

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Me

Me +

25 (2–21%)

(33.14)

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Oxidation of Carboxylic Acids and Derivatives

TAbLE 33.3 Kolbe Cross Coupling Carboxylic Acids

Conditions

EtCO2H + nC5H11CO2H

PrCO2H + nC5H11CO2H

n

C8H17CH — — CH(CH2)nCO2H + RCO2H (1:10) n = 10, R = Et n = 7, R = Me, Pr, nC4H9, nC6H13

aq. Na2SO4 NaOH (Pt) aq. Na2SO4 NaOH (Pt) MeOH NaOMe or KOH (Pt)

Product (Yield, %) nC

4H10 (9)

nC

[84]

7H16 (42),

n n

References

nC

10H22

C6H14, nC8H18 C10H22 (total 27)

nC H CH — CH(CH ) R — 8 17 2n

(28–63)

[84]

[83] [85] [86]

— CHCH2)2(CH2)6CO2H (Z,Z)-nC5H11(CH — + R CO2H MeOH NaOMe (Pt) R = Et, Pr, nC4H9

(Z,Z)-nC5H11(CH — — CHCH2)2(CH2)6R (49–51)

[87]

(Z,Z,Z)-Me(CH — — CHCH2)3(CH2)6CO2H + R CO2H R = Et, Pr, nC4H9

MeOH KOH (Pt)

(Z,Z,Z)-Me(CH — — CHCH2)3(CH2)6R (43–52)

[87]

R1 CO2H + R2OCH2CO2H R1 = nC5H11, R2 = nC6H13 R1 = nC11H23, R2 = nC12H25

DMF KOH (Pt)

R1CH2OR2 (32–58)

EtCH(Me)(CH2)2CO2H + THPO(CH2)7C– – C(CH2)2CO2H

MeOH KOH (Pt)

[88]

EtCH(Me)(CH2)4-C– –C-(CH2)7OTHP [89]

O

O O O

O + R CO2H CO2H

O

O

O R

R = nC7H15, nC11H23 Br(CH2)10

O

(32–33) [90]

O

R1 CO2H + MeO2C(CH2)nCO2H R1 = nC11H23, n = 18 R1 = nC14H29, n = 14 R1 = nC6H13, n = 3, 14

MeOH KOH (Pt) or MeOH NaOMe (Pt)

R1 (CH2)nCO2Me

nC

MeOH/Hex KOMe (Pt) MeOH KOMe

nC

18H37CH(Me)CH2CO2H + MeO2C(CH2)3CO2H

MeOH KOH (Pt)

nC

HO(CH2)18CH(Me)CH2CO2H + MeO2C(CH2)3CO2H (1:6)

MeOH KOH (Pt)

17H35CO2H +

MeO2CCH2CH(Me)CH2CO2H nC

[48] [91] [47] 18H37CH(Me)CH2CO2Me

(25) 18H37CH(Me)(CH2)4CO2Me

[92] [93] [93]

(66) HO(CH2)18CH(Me)(CH2)4CO2Me (31)

[94] (Continued )

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1278

Organic Electrochemistry

TAbLE 33.3 (Continued ) Kolbe Cross Coupling Conditions

Carboxylic Acids

Product (Yield, %)

References

HO(CH2)17CO2H + MeO2CCH2CH(Me)CH2CO2H

MeOH KOH (Pt)

HO(CH2)18CH(Me)CH2CO2Me

[94]

R1R2CH(CH2)mCO2H + RO2C(CH2)nCO2H R1 = Et, nC9H19 R2 = Me, Et, nC5H11 R = Me, Et

MeOH KOH (Pt)

R1R2CH(CH2)m+nCO2R

[95]

MeCO(CH2)2CO2H + MeO2C(CH2)3CO2H

MeOH KOH (Pt)

MeCO(CH2)5CO2Me (56)

[96]

nC H CO H + 8 17 2

MeOH NaOH (Pt)

nC

[97]

— 10H21CH—CH(CH2)2CO3Me

—CH(CH2)2CO2H MeO2C(CH2)2CH—

(48)

—CH(CH2)mCO2H + RCH — MeO2C(CH2)nCO2H R = H, Et, nC4H9 m = 1,2,8 n = 2,3,7,8

MeOH NaOMe (Pt) MeOH KOMe (Pt) MeOH KOH (Pt)

RCH— —CH(CH2)m+nCO2Me (33–97)

[98] [99] [90] [100] [99]

R C– –C (CH2)3CO2H + MeO2C(CH2)nCO2H R = Me, Et, Pr, n = 2,3,4,6

MeOH KOH (Pt)

R C– – C (CH2)n+3CO2Me (49–59)

[101]

MeO2C MeO2C Dn Ph

MeO2C

CO2H +

CO2H MeO2C

CO2H + Dn CO H 2

nC H CH(Me)(CH ) CO H + 18 37 24 2 MeCOCH(Me)(CH2)3CO2H

tBu

DnDn MeOH/Py NaOH (Pt) MeOH KOH (Pt)

nC

18H37CH(Me)(CH2)7CH(Me)(CH2)3COMe

(43)

MeO2C(CH2)5

CO2H N H + Me(CH2)nCO2H

[98]

NHBoc CO2tBu

[104]

(25)

+MeO2C(CH2)3CO2H R1N(Ac)(CH2)2 CO2H +MeO2C(CH2)nCO2H R1=nC12H25, n = 3 R1=nC12H25, n = 4 R1=nC18H37, n = 4

[102] [103]

CO2Me

Ph

[MeOH]/Py–NaOMe–(Pt)

CO2

[97]

(9)

CO2H

BocNH

CO2Me

MeOH/Py–NaOMe–(Pt)

R1N(Ac)(CH2)n+2CO2Me (56–60)

[105]

(CH2)nMe MeOH–NaOH–(Pt/Ti/C)

[106] N H

(Continued )

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Oxidation of Carboxylic Acids and Derivatives

TAbLE 33.3 (Continued ) Kolbe Cross Coupling Carboxylic Acids

Conditions

Product (Yield, %)

References

Br(CH2)10CO2H + MeO2C(CH2)7CO2H

MeOH–KOH–(Pt)

Br(CH2)17CO2Me (76)

[99]

HOCH2CH(CF3)CO2H + R(CH2)2CO2H R = Me, OH, Cl

MeOH–NaOMe–(Pt)

HOH2CH(CF3)(CH2)2R (14–70)

[100]

nC F (CH ) CO Me (68) 8 17 29 2

[107]

nC F (CH ) CO H + 8 17 22 2 MeO2C(CH2)7CO2H

CF3CO2H + + EtO2CCH2CO2H

H2O–KOH–(Pt)

CF3CH2CO2Et (48)

[108]

CF3CO2H + [CD3CO2H]

CD3CO2K–(Pt)

CF3CD3 (68)

[109]

F F F O

MeCN/MeOH/Py–(C)

(3:1) F F F CF3 nC F O CO2H O 3 7 F CF3 ClF2C

O F F3C F F F O O C3F7n F3C F F F

CO2H +

[110]

CO2H +

F Cl

F3C F

MeOH–KOH–(Pt) ClF2C

F 3C F CF3O(CF2)3O

O

[111]

(CF2)3OCF3

F Cl

CO2H

(Z)-n-Enoic acids (27) (n > 5) fully retain the alkene configuration in the Kolbe electrolysis to give the cross coupling products 28, but cyclic products (29) are obtained, to some extent, for 6- and 7-alkenoic acids (27, n = 4 and 5) (Equation 33.15). CO2H

n

+

MeCO2H

MeOH–KOH–(Pt)

27 Me Me

+

(CH2)n

n

28

C.

27 n 3 4 5 6 7

28 68 46 65 72 73

Yield, % 29 — 21 5 — —

(33.15)

29

DECARBOXYLATIVE RADICAL ADDITION

Electrolysis of carboxylates RCO2− in the presence of alkenes (30) affords the radical addition products 31–34. Plausible reaction pathways are illustrated in Equation 33.16. The radical R• generated by electrodecarboxylation of RCO2− first combines with the alkene (30) to give the radical intermediates (35), which may provide the dimers (31) or the radical coupling products (32). Further oneelectron oxidation of 35 may provide cations (36), which subsequently react with the nucleophiles Nu− or liberate H+ to give substituted products (33) and/or alkenes (34).

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1280

Organic Electrochemistry

TAbLE 33.4 Kolbe Cross Coupling of Chiral Carboxylic Acids Conditions

Carboxylic Acids

Product (Yield, %)

tBu

References

tBu

O

O

O

O

MeOH–Et3N–(Pt) OH

O + RCO2H

R O

R = Me2CH (35) (CH2)3NHC(O)H (51) (57) Et Intermediate for macrolides

HO2C

O + RCO2H (1:8)

O

[114,115]

O

MeOH–Et3N–(Pt)

O

R

O

[116,117]

O

R = nC12H25 (40) R = Me(OCH2CH2O)(CH2)12 (40) R = nC4H9 (40) CO2H

O

O

MeOH–KOH–(Pt)

OAc + PrCO2H (1:10)

Pr

O

[118]

OAc (10) O

O CO2H

n CO2R

MeOH–NaOH–(Pt)

[119,120]

n = 4–7,R = Me, Et (10–43)

+ RO2C(CH2)nCO2H (4–6 equiv) R1

O

H N

OR2

R1

CO2H CO2H

+ N H

OR2 [104,121,122]

MeOH/Py–NaOMe–(Pt) R3

R3

O

H N

OR4

OR4

N H

O

O

Y –

RCO2

30

–e– –CO2

R

R

Y 35 –e–

+R

Dimerization Y R

R

R R

R

Y

Y

+Nu–

Y 36

32

31

+

–H+ Nu R

R

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Y

Y 33

34

(33.16)

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Oxidation of Carboxylic Acids and Derivatives

Though the radical addition to the alkenes (30) generally gives rise to a mixture of products 31–34, in some particular cases, reasonable yields of the radical adducts (31 and/or 32) were obtained. Some typical examples are shown in Table 33.5 [107,108,124–137]. Both styrene and conjugated dienes (30, Y = phenyl, alkenyl) afford the corresponding radical adducts 31 in fairly good yields. Addition of the electrogenerated trifluoromethyl radical CF3• to alkenes has been found to be a good procedure for the preparation of trifluoromethylated products 31–34 (Y = CF 3). The electrolysis of partially neutralized trifluoroacetic acid and alkenes (30) generally gives a mixture of mono- and bis-trifluoromethylated derivatives, types 31 and 32 (R = CF 3), and it is quite hard in some cases to separate the products in a pure state. Recently, however, several devices have been made to overcome such difficulties. ω-Trifluoromethylalkanoic acids (37 and 38) [128,129], β-trifuoromethylaldehyde (39) [130], α-trifluoromethylketones (40) [131], and 5-trifluoromethyluracil (41) [133] have thus been synthesized efficiently (see Table 33.5). Addition of trifluoromethyl radical to aromatic compounds gives a mixture of o-, m-, and p-trifluoromethyl-substituted aromatic compounds in moderate yields [138,139]. Some examples are shown in Table 33.6. Intramolecular radical addition may furnish a simple and straightforward access to cyclic compounds through one-electron oxidation of unsaturated carboxylic acid 42, decarboxylation, radical cyclization, and finally cross coupling with radical R• derived from electrooxidative decarboxylation of 43 (Equation 33.17). Representative examples are shown in Table 33.7 [95,107,140–146]. –e–, –CO2

CO2– + RCO2– X 42

X

43

R

+ R

+ R

MeOH–KOH

X

X

(33.17)

44

X = O, NC(O)Me, NC(O)H R = Me, Pent, MeO2C(CH2)4, etc.

A stereoselective synthesis of prostaglandin precursors (47) has been performed, for instance, by coelectrolysis of 45 and 46 (Equation 33.18) [95,147]. R1O

R2

CO–2 + O

OEt

45, R1= H, Ac

R2CO–2

–e– –CO2

R1O

(33.18) O 47 (33–35%)

46, R2 = Me, MeO2C(CH2)2

OEt

Radical tandem cyclization to tricyclic compounds 49 is initiated by the Kolbe electrolysis of unsaturated carboxylic acids 48. Successive intramolecular radical addition followed by mixed coupling with a radical R from the Kolbe decarboxylation of coacid MeCO2H (Equation 33.19) [140]. In this radical tandem cyclization, three C–C bonds are formed in either regio- or stereoselective manner. O

O MeCO2H–MeOH–(Pt)

n

m

CO2H

48 (n = 1–3; m = 1, 2)

n

(33.19)

40°C 25 mA/cm2, 1.2 F m

(34–42%)

49

The cross coupling of the electrogenerated trifluoromethyl radical with active methylene compounds may open another access to trifluoromethylated products. Electrolysis of octyl acetoacetate (50)

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Organic Electrochemistry

TAbLE 33.5 Decarboxylative Radical Addition to Alkenes Carboxylic Acid

Alkene

EtO2C−CO2H EtO2C−CO2H MeO2CCH2CO2H MeO2CCH2CO2H MeO2C(CH2)4CO2H

CH2═CHCH═CH2 CH2═CMe2 CH2═CHPh CH2═CMe(Ph) CH2═CHCH═CH2

MeCO2H CF3CO2H CF3CO2H CF3CO2H CF3CO2H CF3CO2H CF3CO2H CF3CO2H

CH2═C(Me)CHO CH2═CH(CH2)6CO2H CH═CHCO2H CH2═CHCH2OH CH2═C(Me)OAc CH2═C(Me)CH2CN CH2═CHCO2R R ═ Me, Et CH2═CHCONH2

CF3CO2H

CH═CHCN

CF3CO2H

CH2═CHSO2R R ═ Ph, Et

CF3CO2H

EtO2CCH═CHCO2Et

CF3CO2H

Product (yield, %)

References

[EtO2CCH2CH═CHCH2−]2 (66) (70) EtO2CCH2CMe2(CO2Et) (9) [MeO2CCH2CH2CH(Ph)−]2 (38) [MeO2CCH2CMe(Ph)CH2−]2 (50) [MeO2C(CH2)5CH═CHCH2−]2 (47) + [MeO2C(CH2)5CH═]2 (47) [EtC(Me)(CHO)−]2 (80) CF3(CH2)8CO2Ha (16) 37 CF3(CH2)2CO2Ha (8) 38 CF3(CH2)2CHOb (25) 39 CF3CH2C(O)Mec (17) 40 CF3CH2C(Me)CH2CN (good) [CF3CH2CH(CO2R)−]2 (24–50) CF3CH2CH(CF3)CONH2 (37) + CF3CH═CHCONH2 (9) CF3CH2CH(CF3)CN (5) + [CF3CH2CH(CN)−]2 (30) CF3CH2CH(CF3)SO2R (24, 25) + [CF3CH2CH(SO2R)−]2 (22, 22) [EtO2CCH(CF3)−]2 (47) + [EtO2CCH(CF3)CH(CO2Et)−]2 (27)

O

[124] [107] [107] [124] [107] [125,126] [127] [128] [129] [130] [131] [132] [108,133–136] [136] [134,135] [136] [133] [133]

O CF3 N Et

N Et (44) CF3

O CF3CO2H CF3CO2H

O

HC≡C−C4H9

n

O H O

CO2H

O H

N N Me

CF3

N

O

CH2═CHCO2Me

F10

[134,135] [133]

CF3CH═C(CF3)C4H9n

N Me

(40) 41

[137]

CH2— —CHCO2Me F10 CH2CH2CO2Me F10 CH2CH(CO2Me)F10

a b c

After hydrogenation (Pd/C). After hydrogenation and hydrolysis (aq. NaOH). After hydrolysis (aq. H2SO4).

© 2016 by Taylor & Francis Group, LLC

2

1283

Oxidation of Carboxylic Acids and Derivatives

TAbLE 33.6 Decarboxylative Radical Addition to Aromatic Compounds Carboxylic Acid

Aromatic Compound

CF3CO2–

CN

Product (Yield, %)

MeCN/TFA-Pyridine-(Pt)

CF3

References (65)

CN

[138]

o/m/p/ = 46/18/36 –

CF3CO2

COCH3

CF3CO2–

CHO

CF3

COCH3

(40)

[138]

CHO

(52)

[138]

(7)

[139]

CF3

CF3CO2H

CF3

N

N α/β/γ = 27/38/35

in an MeCN–H2O–CF3CO2H–(Pt) system gives a mixture of octyl 2-trifluoromethylacetoacetate (51) and octyl 3,3,3-trifluoropropionate (52), the latter of which is probably produced by deacetylation of 51, with a base electrogenerated at the cathode (Equation 33.20) [147,148]. O

O n

O–C8H17

+ CF3CO2H

–e– –CO2

O

O O–C8H17n CF3 51

50

F3C

D.

(33.20)

O

EG base

O–C8H17n 52

MISCELLANEOUS

Further applications of the radical-induced reactions in the Kolbe decarboxylation reactions involve hydrogen abstraction and coupling with in  situ generated radical species from additives. For example, hydrogen abstraction by radical intermediates derived from paraconic acids (53) occurs selectively in MeOH–NaOMe–Fe powder–(Pt) and MeOH–NaOMe–(C) systems to give lactone 54 (Equation 33.21) [10]. The effect of Fe powder is significant since butenolides (55) are obtained exclusively in an MeOH–NaOMe–(Pt) system. R1 –e– R1

CO2H

R2

R2 O

MeOH–NaOMe–Fe–(Pt) or MeOH–NaOMe–(C)

O

O

54 (92–99%)

R1 O 53

–2e– MeOH–NaOMe–(Pt)

R2 O 55 (68–90%) O

© 2016 by Taylor & Francis Group, LLC

(33.21)

1284

Organic Electrochemistry

TAbLE 33.7 Decarboxylative Radical Cyclization Unsaturated Carboxylic Acid

Carboxylic Acid

Product (Yield, %)

References

Synthesis of Cyclopentane Derivatives (CH2)4CO2Me CO2H

R1

CN CO2Et CN

MeO2C(CH2)4CO2H R

R2 R2 O

2

R2

O

CH3CO2H MeOH–(Pt)

R2 =

R1 =

R1

Me Me H

(75) (75) (71)

[107]

[140]

CO2H

(72) R

EWG RCO2H MeOH–KOH–(Pt)

HO2C O

EWG [141]

O (65–90)

O

O

Synthesis of Oxygen-Containing Cyclic Compounds CO2–

CH3CO2H

[107] [142]

(72) O CH2CO2Me

O CO–2

MeO2CCH2CO2H

[107] (43) O

O O

AcO

O MeCO2H

AcO

O

AcO

O [107]

AcO

C O 2H

(50) R1

CO2H

AcO

R

2H

OEt

O

[143,144]

R

CO2Et

CO2Et

RCO2H MeOH–KOH–(Pt) O

O CO2H CO2Et

CO2Et

R2 R1

[141]

R = Me, Et, MeO2C(CH2)2, — CH2—CHCH 2, etc. (72–90)

R2CO2H

OEt

O

(35) (33)

OEt

O

HO2C

R1

R1 = Me MeCO(CH2)2

AcO 1CO

O

R1 = H, Me R2 = Me, nC5H11 OEt (45–70)

[145]

Synthesis of Nitrogen-Containing Cyclic Compounds CO–2 n

Hex

(67)

C5H11CO2H

N CHO

[107] [95]

N CHO

(Continued)

© 2016 by Taylor & Francis Group, LLC

1285

Oxidation of Carboxylic Acids and Derivatives

TAbLE 33.7 (Continued ) Decarboxylative Radical Cyclization Unsaturated Carboxylic Acid

Carboxylic Acid

Product (Yield, %)

CO2H

R1 R1CO2H N Ac

N Ac CO2H

R1 R1CO2H

N CHO CO2H

N CHO

R1 = Me nC H 5 11 MeC(O)(CH2)4

(58) (46) (53)

[146]

R1 = Me nC H 5 11

(58) (46)

[146]

(56)

[146]

(67) (63)

[146]

Me

Me +

MeCO2H

N Ac

References

N Ac

CO2H

AcO (4:1)

N Ac

R1 R1CO2H

N CHO

N CHO

R1 = Me nC H 5 11

CO2H (45)

MeCO2H

[146]

N CHO

N CHO

Electrolysis of carboxylic acids in the presence of nitrate [149,150], difluoroamide [151], azide [152], and bromide ions [153] has been examined in which radical species generated from both carboxylic acids and additive anions are formally coupled to give substituted products (Equation 33.22). –e–

Y– RCO–2

–e–

Y R–Y Y–: NO–3, F2N–, N–3, Br–

R

–CO2

(33.22)

Substitution of an aldehyde hydrogen (R-CHO) with methyl radical has been performed by electrolysis of acetic acid in the presence of aldehydes yielding methyl ketones (Equation 33.23) [154,155]. –e– –CO2

PhCOMe

6H13–CHO + MeCO2H

–e– –CO2

nC H COMe 6 13

nC

Ph–CHO + MeCO2H

(33.23)

Initiation of polymerization of alkenes with radicals formed in the Kolbe electrolysis has been studied to some extent [156–160].

IV. CARbENIUM ION REACTION (TWO-ELECTRON OXIDATION) The two-electron process in the electrodecarboxylation of carboxylates has been intensively investigated from the mechanistic and synthetic viewpoints. Significant progress of such carbenium reactions has been recognized in alkoxylation, acetoxylation, hydroxylation, acetamidation,

© 2016 by Taylor & Francis Group, LLC

1286

Organic Electrochemistry

alkene formation (β-elimination), rearrangement, C–C bond cleavage, and so on. Typical reaction patterns of the electrogenerated carbenium intermediates R+ and their synthetic potentialities are discussed next.

A. ACYLOXYLATION, ALKOXYLATION, HYDROXYLATION, AND ACETAMIDATION Sufficiently stable electrogenerated carbenium ions R+ can be trapped with nucleophiles or nucleophilic solvents. Acetoxylation and methoxylation proceed by electrolysis in either acetic acid or methanol. Acetamidation occurs in wet acetonitrile, in which the nucleophilic attack of the nitrile group on the carbenium ions R+ tends to give iminium cations 56; subsequent hydrolysis with water in the medium completes the reaction (Equation 33.24) [161]. R΄–OH (R΄ = Ac, Me, H, etc.) –2e– –CO2

R–CO– 2

R–OR΄

R+ +

R–N — — C–Me 56

MeCN

(33.24)

H N

H2O

Me

R O

Electrogenerated carbenium ions, which are stabilized by neighboring heteroatoms, that is, oxygen, nitrogen, sulfur, and others, can lead smoothly to the corresponding acetoxylated and/or methoxylated product (Equation 33.25). This method plays an important role in various synthetic applications. Some typical examples are shown in Tables 33.8 and 33.9 [107,162–189].

R1

R2

R

R2

R1

R

R

–2e– –CO2

R2

R1

Y–

+

N

R

Y–

R2

R2N Y R1

R2 S+

RO Y R1

R2

R1

–CO2

RS CO2H

O+

R1

–2e–

R2N CO2H 1

R

–2e– –CO2

RO CO2H

Y–

RS Y R1

R2

(33.25)

R2

R2

Y = AcO, MeO, etc.

α-O-Substituted carboxylic acids can give mixed acetals; the electrolysis of uronic acid derivative (57) in methanol thus affords the mixed acetals (58) (Equation 33.26). The method has been successfully applied to aminocyclitol synthesis as well as to the selective fission of glucuronide linkages in oligoglycosides (59) (Equation 33.27) [190]. CO2H O OH

MeO OH

O

MeOH–Et2NH–(C)

OH HO

HO OH 57

© 2016 by Taylor & Francis Group, LLC

1) MeNO2–NaOMe 2) Ra–Ni OH 3) Ac2O

OH 58 (51%)

MeO R OH

OH

HO OH

R = NO2, NHAc

(33.26)

1287

Oxidation of Carboxylic Acids and Derivatives

TAbLE 33.8 Acyloxylation, Alkoxylation, and Hydroxylation of α-O-Substituted Carboxylic Acidsa Conditions

Carboxylic Acid CO2H O

Products (Yield, %) OAc O

AcOH–Et3N–(Pt)

OMEM

OMe HO2C AcNH O OMe OMe

MeOH–NaClO4–(Pt)

OMe MeO AcNH O OMe OMe

(96)

[163]

MeO

MeO OMe MeO OMe O O

O OMe MeO CO2H

OMe MeO OMe O

MeOH–NaClO4–(Pt)

CO2H

O

O

HO

(93)

[107]

(92)

[107]

OMe O

O O MeO

MeOH–NaClO4–(Pt)

OMe OMe

O O

O O

HO

CO2H R

R

MeOH–NaOMe–(C)

CO2H HC OH HO CH HC OH HC OH CH2OH CO2H R O R = CH2CH — —CH2 CH2C(Me) — — CH2, etc R2 R3

O 1

R CO2H O R1 = Ph, Me, EtO R2, R3 = H, Me, Ph,–(CH2)5–

R1= OH H H H H2O–NaOH–(Pt) flow cell

R4= H H H OH

(93) (90) (98) (87) [165–167]

HO OH

MeOH–K2CO3– Et4NClO4–(Pt)

MeOH–(C) or MeCN–Et3N–(C)

O

O R1

(67–90)

OMe R

R2 R3 O

[168]

[169,170]

OMe

Me MeCN–(C)

MEM = β-methoxyethoxymethyl.

© 2016 by Taylor & Francis Group, LLC

[164]

O OH HO

Me CO–2

OMe R3

R4 R3= H H OH H

R2 R2= H OH H H

R4

R2

O

R1

3

1

a

[162]

(89)

OMEM

MEMO

MEMO

EtO

References

O

EtO O

OEt Me

(81)

[171]

1288

Organic Electrochemistry

TAbLE 33.9 Acyloxylation, Alkoxylation, and Hydroxylation of α-N-Substituted Carboxylic Acids Carboxylic Acid

Conditions

Products (Yield, %)

HO2C

N COR1 R = H, TBDMS R1= Me, OMe O O

HN HO2C

References

OR

OR MeOH–Et3N–(Pt)

MeO

(89–97)

N COR1

O

Me

O O

HN

MeOH–NaOAc

[172,173]

MeO

O

HN

[174,175]

Me

Me MeO

(95,77/23) O

O CO2H N CO2Me R = iPr, tBu

R

N Ac

O

Ph CO2Et CO2H

CO2H

N R

AcOH

R

CO2Me (85–87) MeOH–NaOMe–(C) N Ac

MeOH–NaOMe–(C)

O

N

O

O

[177,178] (98)

OMe (75–98)

[179]

tBu

CO2Me CO2H O

N

Ph CO2Et OMe

N R

R = H, nC4H9

tBu

HN

[176]

OC(O)Me

N

HN MeOH–NaOAc–(C)

CO2H

O

NH R = H, Et

AcNH CO2Et CO2–

OMe

R R = H, alkyl, allyl, etc

(70)

OAc

R

NH

(76)

O R΄OH–(C) R΄= Me, Et, iPr

[181]

NH

O

MeCN–AcOH–NaOAc –(Pt)

[180]

O

N

CO2H R

CO2Me OMe

O

Et3N/MeOH–(C)

NH

N

AcNH R

[182]

CO2Et OR΄

[183,184] (Continued )

© 2016 by Taylor & Francis Group, LLC

1289

Oxidation of Carboxylic Acids and Derivatives

TAbLE 33.9 (Continued) Acyloxylation, Alkoxylation, and Hydroxylation of α-N-Substituted Carboxylic Acids Carboxylic Acid

Conditions

Products (Yield, %)

MeOH–NaOMe–(C)

CO2H

OMe

EtOH–NaOEt (0.2 equiv.)– (Pt)

NHBoc CO2H

(79)

[186]

(63, 70)

[186]

O ROH–NaOR (0.2 equiv.)– (Pt)

O

HN

CO2H Me

R = Me, Et

MeOC(O)

HN Me

O

OR Me

MeOC(O)

N

CO2H

O

N

MeOH–NaOMe–(Pt)

OMe O

[187]

O

O R1

N

O

[185]

NHBoc OEt

O

H

(95, 96)

NHAc

NHAc

Me

References

2OH–DBU–(Pt)

R R2 = Me, CD3, Et, Ac, d3-Ac

N CH2 CO2H

H

R1

N

(35, 58)

[188]

R = Me (98)

[189]

N CH2

O

OR2

R1 = H, Me H Z

Z

N H N

N H

MeOH, 2.3 F AcOH, 2.3 F

CO2H

H Z

O

Z

BzIO

iPr

O

N H

CO2H

H N

O

MeOH, 2.2 F

Z

N

H N

OSiMe3 CO2H NHCO2Me

© 2016 by Taylor & Francis Group, LLC

Z

MeOH, 2.4 F

BzIO MeOH–Et3N–(Pt)

N H

OR

N H

OMe

H N

O

R = Ac (67)

O

CO2H N

iPr

O

(98)

[189]

(50)

[189]

(75)

[176]

OMe N

OSiMe3 OMe NHCO2Me

1290

Organic Electrochemistry OH HO HO O

HO HO

Me HO HO O O Me

Me

OH O O OH

1) –e–, AcOH–Et3–(Pt) 2) MeNO2/NaOMe–MeOH

HO

HO

OH

HO

O O

HO

HO

OO

HO

OH OH

CO2H O O O

O OH

O R

O

OH O O OH

60 (R = NO2, 58%)

HO

HO O O Me

OH

(33.27)

HO OH +

OH 59

O

61a: R = α-OH, β-H 61b: R = O (16%)

R HO

HO OH

Electrolysis of the dicarboxylic acid (62) in an MeOH–MeONa–(Pt) system and subsequent workup afford the keto acid (63), a precursor of chrysanthemic acid, in 86% yield via intramolecular acyloxylation and acid-promoted hydrolysis (Equation 33.28) [191].

OMe

OMe

CO2H MeOH–MeONa–(Pt) CO2H 86%

H+

O

O CO2H

O

62

(33.28)

63

Decarboxylative acetoxylation and methoxylation of α-aminocarboxylic acids proceed smoothly, yielding synthetically useful intermediates. For example, the versatile intermediate (65) for the synthesis of thienamycin is prepared by electrodecarboxylative acetoxylation of 4-carboxy-2-azetidinone (64) in an AcOH–MeCN–NaOAc–(Pt) system (Equation 33.29; TBDMS = tert-butyldimethylsilyl) [182,192]. TBDMSO

H H CO2H NH

O 64

TBDMSO MeCN/AcOH–NaOAc–(Pt)

H H OAc NH

84%

(33.29)

O 65

Electrolysis of N-carbamoylaspartic acid or N-ethoxycarbamoylaspargine in MeOH– NaOMe–(C) affords the corresponding methoxylated products [193]. A 5-fluorouracil derivative (68), a potent antitumor agent, is prepared via electrolytic methoxylation of N-acylazacycloalkane-2-carboxylic acids (66) in MeOH–NaOMe–(C) and subsequent condensation of (67) with 2,4-bis(trimethylsilyl)-5-fluorouracil (TMS-5-FU) (Equation 33.30) [179,194]. Recently, the decarboxylative methoxylation of α-amino acids has been extended to prepare useful chiral building blocks [172–174,176].

© 2016 by Taylor & Francis Group, LLC

1291

Oxidation of Carboxylic Acids and Derivatives O F

HN

O

O CO2H

N O

n

MeOH–NaOMe–(C)

n

N SiMe3

80–95%

OMe

N O

R 66 n = 1,2 R = H, Me, CH2NHCO2Bn

F

HN O

SnCl4

R

n

(33.30)

O

N N

R

68

67

The decarboxylative methoxylation of N-benzoyloxazoline derivative 69 derived from L-serine in MeOH–NaOMe–(C) gives optically active N,O-acetal 70 (39% ee) as an inversion product (Equation 33.31) [195]. This is the first example of memory of chirality through an electrogenerated achiral carbenium ion. Among the anode materials examined, only graphite gives positive results concerning the ee% of the product, while other materials, such as glassy carbon, Pt, and Au, give racemic products. O Me Me

CO2H

N COPh

O

–2e– MeOH–NaOMe–(C) 25 mA/cm2, –20°C

Me Me

OMe N COPh

(33.31)

70 (69%, 39%ee)

69

Electrolysis of α-sulfenyl carboxylic acids in protic solvents gives ketones, presumably via replacement of the carboxyl group with alkoxyl or acetoxyl groups and subsequent one-electron oxidation of the sulfenyl moiety (Equation 33.32) [196]. (This discussion is postponed to Section IV.D).

RS CO2H 1

R

R2

–2e–, –CO2 R΄ OH

RS OR΄ R1

R2

O

–e– –RS

R1

R2

(33.32)

Further selected examples of the decarboxylative acetoxylation and methoxylation are listed in Table 33.10 [26,197–206]. The electrogenerated stable carbenium ions, such as benzylic, allylic, homoallylic, and tertiary carbocation intermediates, are likely to be trapped by suitable nucleophiles, yielding the corresponding acetates, methyl ethers, and, in some cases, intramolecular acyloxylation products.

B. ALKENE FORMATION The electrodecarboxylation becomes a good procedure for the formation of carbon–carbon double bonds, which involves roughly two types of reactions: (1) decarboxylative discharge of carboxylates followed by releasing a vicinal leaving group (Equation 33.33a) and (2) initial discharge of hetero atoms at the α- or β-position to carboxylic acid, which causes an elimination of the carboxyl group in a concerted manner (Equations 33.33b and c).

© 2016 by Taylor & Francis Group, LLC

1292

Organic Electrochemistry

TAbLE 33.10 Electroacyloxylation and Alkoxylation Carboxylic Acid

Conditions

R CO2H

Products (Yield, %)

AcOH/AcOEt/tBuOH–Et3N–(Pt)

References [197,198]

OAc

CO2–

MeOH

CH2CO2H

n

[199]

OMe (100) CH2OMe

n

[202]

MeOH–NaOMe–(C)

CO2Et

CO2Et

n = 1,2

(35, 28) CO2H

OAc AcOH/tBuOH–Et3N–(Pt)

[203,204]

O

tBu

O

(72)

tBu

CH2CO2H

tBu-CO

MeOH–NaOMe–CO2–(Pt)

2H

CH2OMe

[26,199,201]

[205]

tBu-OCO Me 2

(78) CO2H O MeOH–(Pt)

[206] (77)

CO2H

Y C C CO2–

Y C

–2e– –CO2 (Type l–IV)

H X C C

–e–

C +

C C

(33.33a)

Y = H, CO2H, SiMe3, CH2OH, etc. H X C C

+

CO2H

CO2H –e–

–CO2

C C

–X (Type V)

–H+

+

X C C CO2H

© 2016 by Taylor & Francis Group, LLC

–Y+

O

(33.33b) –CO2

X

–H+ (Type VI)

C C

1293

Oxidation of Carboxylic Acids and Derivatives +

H X C C

–e–

CO2H

H X C C CO2H – –CO2, –H+ –e

X

–H+

–

– – –

H X+

(33.33c)

C C

C C (Type VII)

Various sets of functional groups have been devised for both categories of decarboxylative carbon– carbon double-bond formation. Representative examples of a set of functional groups together with their electrolysis conditions are summarized in Table 33.11 [11,119–200,207–214]. In such alkene formation reactions, an electrodecarboxylation–deprotonation sequence and double decarboxylation (Types I and II) are involved. Electrochemical aromatization of the nonconjugated dienes (71 and 73) can be successfully performed by this method as in Equations 33.34 and 33.35 [215]. Decarboxylation of the γ-lactone α- and β-carboxylic acids (75 and 77) in a Py–H 2O–Et3N–(Pt or C) proceeds smoothly, yielding the corresponding butenolides (76 and 78), respectively, as in Equations 33.36 and 33.37 [216,217]. In a similar fashion, 2-substituted cyclopentenones (80) have been obtained starting from 2-carboxycyclopentanones (79) (Equation 33.38) [218].

nC

5H11

n

CO2H

C5H11

MeOH–NaOMe–(Pt) MeO

(70%) MeO

OMe

(33.34)

OMe

71

72 Me

Me CO2H MeOH–NaOMe–(Pt)

(65%)

73

(33.35)

74

CO2H Py/H2O–Et3N–(Pt)

O O

O

75

Py/H2O–Et3N–(Pt)

HO

H

(57–63%)

HO

H

Et

H2O–KOH–(Pt)

79

© 2016 by Taylor & Francis Group, LLC

O

Et O 80

(33.37)

O

78

O

CO2H

(33.36)

76

CO2H

O

77

O

(95%)

O

(80%)

(33.38)

1294

Organic Electrochemistry

TAbLE 33.11 Decarboxylative Formation of Carbon–Carbon Double bond Type

Set of Functional Groups

Conditions

References

Direct Discharge of Carboxylates

I

H C C CO2–

Py/H2O Et3N (Pt)

[11]

Py/H2O Et3N (C)

[199,200] [207–209]

MeCN/EtOH KOH (C)

[210]

MeOH Et3N (Pt)

[207] [211]

–2e– anode

II

CO2H C C CO2–

–2e– anode

III

SiMe3 C C –2e–

CO–2

anode

IV

C O H C C CO–2

–2e– anode Concerted Decarboxylation

–e– V

anode

SR C C CO2–

H2O Et3N/LiClO4 (Pt)

[212]

N C C H CO2–

H2O KOH (C)

[225]

MeCN Et4NBF4 (Pt)

[213]

–2e– VI

anode

–2e– VII

© 2016 by Taylor & Francis Group, LLC

anode

N H C C CO–2

1295

Oxidation of Carboxylic Acids and Derivatives

Double decarboxylation of 1,2-dicarboxylic acids (Type II) may take place smoothly in an aqueous 90% pyridine–Et3N–(Pt) electrolysis system to give alkenes in fair yields [219–223]. Some representative examples are shown in Equations 33.39 [207], 33.40 [199,208,209], 33.41 [221], 33.42 [222], and 33.43 [223]. MeO2C MeO2C

MeO2C MeO2C

CO2H CO2H CO2Me

(65%)

(33.39)

CO2Me Py/H2O–Et3N–(Pt)

O

(67%)

(33.40)

MeO2C

MeO2C O

O

CO2H CO2H

Py/H2O–Et3N–(Pt)

(45%)

(33.41)

O O

F

F MeCN/Py/H2O–Et3N–silica gel–tbutylcatechol–(Pt)

O F

H H

(33.42)

F

(17%)

CO2H CO2H Py/H2O–Et3N–(Pt)

(10%)

(33.43)

O O

O

(42%)

O

A new device for alkene formation via decarboxylation–desilylation (Type III) has been achieved successfully by electrolysis of β-trimethylsilyl carboxylic acids in an MeCN–EtOH–KOH–(C) system [210]. For example, norbornadiene (82) and cyclohexadiene (84) have been obtained in good yields (Equations 33.44 and 33.45). SiMe3

MeCN/EtOH–KOH–(C)

(33.44) 82 (76%)

CO2H 81 SiMe3 MeCN/EtOH–KOH–(C)

(33.45)

CO2H 83

84 (83%)

Equation 33.46 demonstrates that the double decarboxylation of the dicarboxylic acid (85) preferentially proceeds rather than the decarboxylation–desilylation reaction to give the alkene (86) in 73% yield [224].

© 2016 by Taylor & Francis Group, LLC

1296

Organic Electrochemistry CO2H CO2H

Py/aq. KOH (0.3 M)–(C)

(33.46) O

SiMe3

O

SiMe3 86 (73%)

85

Electrolysis of γ-hydroxycarboxylic acid gives rise to decarboxylative β,γ-bond fission, leading to olefinic ketones (Type IV). An example is given in Equation 33.47 [211]. CO2H H

(33.47)

O

34–38%

O

H

H

The α-exo-methylene-γ-lactone framework has been successfully constructed via two electrosynthetic pathways, that is, both by the direct (Type I) and by the concerted decarboxylation processes (Type V) as mentioned earlier (Equation 33.48) [212]. The electrodecarboxylation of 87a is probably initiated by a one-electron oxidation of the sulfur atom, giving first the cation radical (87b) and subsequently a concerted elimination of the thiyl radical and carbon dioxide to α-exomethylene-γ-lactone 88 (Type V). On the other hand, the electrochemical decarboxylation of 89a involves an E1-type elimination of a proton from the cation intermediate (89b) generated from direct two-electron oxidation of the carboxyl group (Type I). The latter method generally requires a higher oxidation potential than that required for the concerted method. Therefore, the concerted electrodecarboxylation method becomes more advantageous, especially when the substrates or products are unstable under oxidative conditions. R1 R2

R3 CO2H

O

–

–e 73–92%

SPh O 87a

R1 R2

R3 –2e– O

R1 R2 O

–H+, –CO2 –PhS

R1 R2

R3 CO2H

O

CO2H

30–35% O

O 88

–e–

R3

–2e– –CO2

–H+ R1 R2

89a

(33.48)

R3 H

O

+

O 87b

+

SPh

O 89b

The decarboxylation initiated by the oxidation of a nitrogen atom substituted on the α- and β-positions of carboxylic acids promises versatile synthetic applications for making unsaturated nitrogen-containing heterocycle frameworks (Types VI and VII). Pyrrole [178], uracil [193,213,225,226], thymine [213], 2-pyrazoline [214], β-carboline [227], and dihydroquinoline [228], useful intermediates for drug syntheses, have been prepared by using the electrodecarboxylation procedure. Some typical examples are given in Table 33.12.

C. REARRANGEMENT A carbocation formed by an electrodecarboxylation process is expected to be different in nature from carbocations generated in homogeneous media, for example, by solvolysis or deamination [230,231].

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Oxidation of Carboxylic Acids and Derivatives

TAbLE 33.12 Decarboxylative Alkene Formation of α- and β-Nitrogen-Atom-Substituted Carboxylic Acids Carboxylic Acid

Conditions

Ph CO2Et CO2H

N

Product (Yield, %) Ph

THF/H2O KOH (C)

(94)

O

O

HN

H2O KOH (C) CO2H

N H

HN O

[213] N H

O

(91)

O CO2H

HN

AcOH (C) N H

O

[178]

CO2Et

N O

O

O

References

(50) N H

O

H

H

O

O N

N tBu

CH2Cl2 Bu4NPF6 (C)

N CO2– Bu4N+

Ph

[193] [225] [226]

HN

tBu

[229]

N Ph

O

O

Ph

Ph MeCN Et4NBF4 (Pt)

N N Ph

CO2H

N N Ph

[214] (94)

CO2H NH N H

MeOH/H2O KH2PO4/K2HPO4 (C)

CO2H

R1O R2O

N H

N

[227] (75)

R1O NH CO2H

R3

MeOH NaOMe (C)

R2O

N [228]

R3

It is very likely that the carbocation produced in the electrodecarboxylation should be highly reactive, possessing an unsolvated vacant p-orbital and having no counterion in the near vicinity. Actually, certain carboxylic acids bearing not only a suitably oriented β-substituent but also a heteroatom linked to the β-carbon are liable to undergo a carbocation rearrangement. Some reactions are exemplified as follows.

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Organic Electrochemistry

The electrochemical rearrangement of paraconic acids (90) (Equation 33.49), γ-carboxy-δlactones (92) (Equation 33.50) [232], and β-keto carboxylic acid ethylene acetals (94) (Equation 33.51) [233,234] has been known to proceed smoothly. γ-Keto carboxylic acids (96) and acetals (98) undergo decarboxylative rearrangement to afford esters (97 and 99) as in Equations 33.52 and 33.53 [235,236]. (CH2)n (CH2)n

CO2H

MeOH–Et3N–(Pt)

MeO

O O 90

(33.49)

O

O 91 (60%)

CO2H (CH2)n

MeOH–Et3N–(Pt)

(CH2)n

O

CO2Me

O 93

O

(33.50)

92 R2 O

R3

CO2H

O

R3

MeOH–KOH–(C)

1 O R 94

R1 OMe 95 (40–60%) O

CO2tBu O

R2

MeOH–NaOH–(C)

CO2Me CO2tBu 97 (55%)

CO2H 96 CO2H R O

O

(33.51)

1) MeOH–K2CO3–(C) 2) OH– 3) CH2N2

(33.52)

R CO2Me

(33.53)

99

98

The electrodecarboxylation of β-hydroxycarboxylic acids, easily accessible by Reformatsky reaction on cyclic ketones, furnishes a promising route for the ring enlargement of cyclic ketones; for example, electrolysis of 1-hydroxy-2-cyclotetradecenylacetic acid (100) in an MeOH–KOH–(C) system gives the sp2 carbon migration product (101, 30%) together with the sp3 carbon migration product (102, 5%) (Equation 33.54) [237].

CH2CO2H OH

MeOH–KOH–(C)

+

O

(33.54)

O 100

101 (30%)

102 (5%)

A Wagner–Meerwein rearrangement via a bridged carbenium ion has been observed in the electrodecarboxylation of bicyclo[2.2.1]heptane-2-carboxylic acids (103a–c), leading to the formation

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Oxidation of Carboxylic Acids and Derivatives

of the same products (104 and 105) (Equation 33.55) [238]. This method has been successfully applied to the synthesis of useful intermediates for natural product synthesis. Some examples are given in Equations 33.56 and 33.57 [239–243].

CO2H 103a

MeOH–NaOMe–(Pt or C)

CO2H

CO2H 103b

(33.55)

103c OMe 104

105

MeOC(O) CO2Me CO2H

O R

R

D.

CO2Me CO2H

+

AcOH/tBuOH–Et3N–(C)

Others

(33.56)

OAc (56%)

HO R

1) MeOH–NaOMe-(C) 2) HCIO4

OAc

Several steps

CO2Me

(33.57)

C(O)R (87%)

FORMATION OF ALDEHYDES, KETONES, ACETALS, AND α,β-ENONES

Electrodecarboxylation procedures for preparing aldehydes, ketones, acetals, and enones have been intensively investigated. The following sets of functional groups are found to be electrosynthetically equivalent to carbonyl groups:

OH

O

O

OH +

CO2–

CO2–

O CO–2

O

;

OR

CO2–

SR

; CO2–

CO2–

O

CO2–

As mentioned in Section IV.C, γ- and β-hydroxycarboxylic acids undergo decarboxylative rearrangement to afford ketones. An additional example is given in Equation 33.58, in which the electrodecarboxylation of a γ-hydroxylcarboxylic acid moiety in a bicyclo[4.4.0]decane system (106) gives a 10-membered nonconjugated enone (107) [244]. On the other hand, the electrolysis of a β,γ-epoxy carboxylic acid moiety arranged in the same ring system (108) provides a conjugated enone (109) in good yield (Equation 33.59) [197]. Electrodecarboxylation of α-hydroxy or α-alkoxycarboxylic acids proceeds smoothly, affording aldehyde 111 (Equation 33.60) [245–247] or acetals (see Sections IV.A and IV.B).

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Organic Electrochemistry CO2H

(33.58) OH 106

O 107

CO2H MeOH–Et3N–(Pt–SUS)

O O

O

(33.59)

O O 109 (84%)

O

108 OH

CHO

HO

HO

CO2H

H2O–KH2PO4/NaOH–(C)

(33.60)

HO

HO

111 (60%)

110

α-Sulfenyl carboxylic acids can be transformed into aldehydes and ketones via electrodecarboxylation. For example, the electrolysis of α-phenylthiocarboxylic acid (112) gives aldehyde (113) in an H2O–NaOH–(Pt) system and the corresponding acetal (114) in an MeOH–LiClO4 –(Pt) system (Equation 33.61) [248]. Cyclopentadecanone (116) is prepared by electrolysis of α-phenylthio- and methylthio-substituted cyclopentadecanecarboxylic acids (115) (Equation 33.62) [249]. H2O–NaOH–(Pt) Pr

Pr

H

O 113 (48%)

CO2H

(33.61)

SPh MeOH–LiCIO4–(Pt)

112

OMe

Pr

OMe 114 (72%) (CH2)14

CO2H

EtOH–NaOH–(Pt)

SR 115 (R = Ph, Me)

(CH2)14

O

(33.62)

116 (91–97%)

The α-sulfenyl group can be replaced by a benzothiazolthio group, as demonstrated in the synthesis of the 1,4-diketone (118), a precursor of cis-jasmone (Equation 33.63) [250]. The geminal dicarboxylic acid (119) can also be transformed into 118 by electrochemical decarboxylation as in Equation 33.64 [251], suggesting that the substituted malonic acid group is synthetically equivalent to a carbonyl group. Et CO2H S-BT O

AcOH/tBuOH–Et3N–(C) 61% (BT=2-benzothiazolyl)

(33.63)

Et O

117 Et CO2H CO2H O 119

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MeOH/AcOH–NaOAc–(Pt) 55%

O 118

(33.64)

1301

Oxidation of Carboxylic Acids and Derivatives

E.

MISCELLANEOUS

Further synthetically meaningful reactions induced by electrochemical decarboxylation are briefly discussed in this section. Potent stereoselective routes to methyl cis- and trans-4,4-dimethoxy2-butenoates (121 and 122) and methyl 4,4-dimethoxybutanoate (123) from 2-furoic acid (120) have been achieved mostly by simple selection of an appropriate electrolyte and electrode, that is, MeOH–Et3N–Et4NClO4 –(Pt) for 121; MeOH–Et3N–Et4NBr–(Pt) for 122; and MeOH–Et3N– Et4NBr–(C) for 123, respectively, as shown in Equation 33.65 [252,253]. MeOH–Et3N/Et4NCIO4–(Pt) (MeO)2CH CO2Me 121 CO2Me

MeON–Et3N/Et4NBr–(Pt) O

(33.65)

(MeO)2CH 122

CO2H 120

CO2Me

MeOH–Et3N–(C) (MeO)2CH 123

Electrodecarboxylation followed by α,β-C–C bond cleavage has been investigated for three- and four-membered ring systems. The dichlorocyclopropane carboxylic acid (124) undergoes decarboxylation in either acetic acid or methanol to give the ring-opened acetate (125) or methyl ether (126) in good yields (Equation 33.66) [254]. Ph

AcOH–NaOAc–(Pt) AcO CI

CI 125 (94%)

CI CI

(33.66) Ph

CO2H 124

MeOH–NaOMe–(Pt)

Ph MeO

CI 126 (81%) CI

The double bond in alicyclic β-keto acid enol ethers (127 and 129) can be cleaved by electrodecarboxylation, affording diesters 128 and 130, respectively (Equations 33.67 and 33.68) [255]. CO2H MeOH–LiCIO4–(Pt)

CO2Me

OMe 127 MeOH–LiCIO4/H2SO4–(Pt)

O 129

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(33.67)

128 (59%)

CO2H OMe

CO2Me

CO2Me

(33.68) OMe CO2Me 130 (72%)

1302

Organic Electrochemistry

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34

Heterocyclic Compounds Fructuoso Barba and Belen Batanero

CONTENTS I. II.

Introduction ..........................................................................................................................1311 Electrosynthesis of Six-Membered Heterocyclic Compounds .............................................1311 A. Compounds Containing One Nitrogen Atom in the Ring ............................................1311 1. Pyridine and Derivatives........................................................................................1311 2. Quinoline and Derivatives .................................................................................... 1313 3. Isoquinoline and Derivatives .................................................................................1314 4. Phenanthridine and Derivatives ............................................................................ 1315 B. Compounds Containing Two Nitrogen Atoms in the Ring ......................................... 1315 1. Pyrimidines ........................................................................................................... 1315 2. Phenazines ............................................................................................................ 1315 3. Quinazolines ......................................................................................................... 1315 C. Compounds Containing Three Nitrogen Atoms in the Ring........................................1316 1. Triazines and Benzotriazines.................................................................................1316 D. Compounds Containing One Oxygen Atom in the Ring..............................................1317 1. Pyrones and Derivatives ........................................................................................1317 2. Chromenes .............................................................................................................1317 3. Chromones .............................................................................................................1317 4. Coumarins and Derivatives ...................................................................................1318 5. Isocoumarins..........................................................................................................1318 E. Compounds Containing One Nitrogen and One Oxygen Atom in the Ring ................1319 1. Oxazines and Derivatives ......................................................................................1319 F. Compounds Containing One/Two Sulfur Atoms in the Ring ......................................1319 1. Thiopyran, Thiochromone, and Derivatives ..........................................................1319 G. Compounds Containing One Nitrogen and One/Two Sulfur Atoms in the Ring.........1319 1. Benzothiazines .......................................................................................................1319 2. Dithiazines .............................................................................................................1319 III. Electrochemistry of Six-Membered Heterocyclic Compounds........................................... 1320 A. Compounds Containing One Nitrogen Atom in the Ring ........................................... 1320 1. Pyridine and Derivatives....................................................................................... 1320 2. Quinolines and Isoquinolines ............................................................................... 1321 B. Compounds Containing Two Nitrogen Atoms in the Ring ......................................... 1322 1. Pyridazine and Derivatives ................................................................................... 1322 2. Pyrazine ................................................................................................................ 1323 3. Pyrimidine and Derivatives .................................................................................. 1323 4. Cinnolines ............................................................................................................. 1323 5. Phthalazines and Derivatives ................................................................................ 1324 6. Quinoxalines ......................................................................................................... 1324 7. Phenanthroline and Derivatives ............................................................................ 1324 C. Compounds Containing Three Nitrogen Atoms in the Ring....................................... 1324 1. Triazines and Derivatives ..................................................................................... 1324

1309

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1310

Organic Electrochemistry

D. Compounds Containing One Oxygen Atom in the Ring............................................. 1325 1. Pyrylium Salts....................................................................................................... 1325 2. Pyrones.................................................................................................................. 1325 3. Chromones ............................................................................................................ 1325 4. Coumarins............................................................................................................. 1325 E. Compounds Containing One/Two Sulfur Atoms in the Ring ..................................... 1326 1. Thiopyrans and Derivatives .................................................................................. 1326 2. Dithiines and Derivatives ..................................................................................... 1326 F. Compounds Containing One Nitrogen and One Sulfur Atom in the Ring ................. 1326 1. Thiazines and Derivatives..................................................................................... 1326 IV. Electrosynthesis of Five-Membered Heterocyclic Compounds .......................................... 1327 A. Compounds Containing One Nitrogen Atom in the Ring ........................................... 1327 1. Pyrrole and Derivatives ........................................................................................ 1327 2. Indoles ................................................................................................................... 1328 3. Carbazoles............................................................................................................. 1330 B. Compounds Containing Two Nitrogen Atoms in the Ring ......................................... 1330 1. Imidazole and Benzo Derivatives ......................................................................... 1330 2. Pyrazole and Benzo Derivatives ........................................................................... 1331 C. Compounds Containing Three Nitrogen Atoms in the Ring....................................... 1332 1. Triazole and Benzo Derivatives ............................................................................ 1332 D. Compounds Containing One Nitrogen and One Oxygen Atom in the Ring ............... 1333 1. Oxazoles................................................................................................................ 1333 2. Isoxazole and Benzo Derivatives .......................................................................... 1333 E. Compounds Containing Two Nitrogen and One Oxygen Atom in the Ring ............... 1334 1. Oxadiazole ............................................................................................................ 1334 F. Compounds Containing One Nitrogen and One Sulfur Atom in the Ring ................. 1334 1. Thiazole, Isothiazole, and Benzo Derivatives ...................................................... 1334 G. Compounds Containing One Oxygen Atom in the Ring............................................. 1335 1. Furan and Benzo Derivatives................................................................................ 1335 H. Compounds Containing One Sulfur Atom in the Ring ............................................... 1336 1. Thiophene and Benzo Derivatives ........................................................................ 1336 I. Compounds Containing One Sulfur and One Oxygen Atom in the Ring ................... 1336 1. Oxathioles ............................................................................................................. 1336 V. Electrochemistry of Five-Membered Heterocyclic Compounds ......................................... 1336 A. Compounds Containing One Nitrogen Atom in the Ring ........................................... 1336 1. Pyrrole and Derivatives ........................................................................................ 1336 2. Indole and Derivatives .......................................................................................... 1337 3. Indolizines............................................................................................................. 1340 B. Compounds Containing Two/Three Nitrogen Atoms in the Ring .............................. 1340 1. Imidazole and Derivatives .................................................................................... 1340 2. Pyrazole and Derivatives ...................................................................................... 1340 3. Triazoles and Benzotriazoles ................................................................................ 1341 C. Compounds Containing One Oxygen/Sulfur Atom in the Ring ................................. 1342 1. Furan and Benzo Derivatives................................................................................ 1342 2. Thiophene and Benzo Derivatives ........................................................................ 1343 D. Compounds Containing One Nitrogen and One Oxygen/Sulfur Atom in the Ring.... 1343 1. Isoxazoles and Benzoxazoles ................................................................................ 1343 2. Thiazoles ............................................................................................................... 1344 E. Compounds Containing Two Nitrogen and One Sulfur Atom in the Ring ................. 1344 1. Thiadiazole and Derivatives ................................................................................. 1344

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Heterocyclic Compounds

1311

VI. Electrosynthesis of Fused Heterocyclic Compounds .......................................................... 1344 A. Compounds Containing Two Nitrogen Atoms in the Molecule .................................. 1344 1. Naphthyridines and Derivatives............................................................................ 1344 2. Quinazolines ......................................................................................................... 1345 3. Pyrrolopyrazine .................................................................................................... 1346 4. Pyridoimidazole .................................................................................................... 1346 B. Compounds Containing Three Nitrogen Atoms in the Molecule ............................... 1346 1. Pyrimidoindol Derivatives .................................................................................... 1346 C. Compounds Containing Three Nitrogen and One Oxygen Atoms in the Molecule ... 1346 1. Imidazo-Oxadiazines............................................................................................ 1346 D. Compounds Containing Three Nitrogen and One Sulfur Atom in the Molecule ....... 1347 1. Thiazolotriazines .................................................................................................. 1347 E. Compounds Containing Four Nitrogen and One Sulfur Atom in the Molecule ......... 1347 1. Triazinothiadiazines ............................................................................................. 1347 VII. Electrochemistry of Fused Heterocyclic Compounds ......................................................... 1348 A. Compounds Containing Three Nitrogen Atoms in the Molecule ............................... 1348 1. Pyrazolopyrimidines ............................................................................................. 1348 B. Compounds Containing Four Nitrogen Atoms in the Molecule ................................. 1348 1. Purines and Derivatives ........................................................................................ 1348 2. Pteridines .............................................................................................................. 1350 C. Compounds Containing Two Nitrogen and One Oxygen Atom in the Molecule ........ 1350 1. Pyridooxazines...................................................................................................... 1350 D. Compounds Containing Two Nitrogen and One Sulfur Atom in the Molecule .......... 1350 1. Thienopyrazines.................................................................................................... 1350 E. Compounds Containing Three Nitrogen and One Oxygen Atom in the Molecule ..... 1350 1. Imidazo-Oxadiazines............................................................................................ 1350 F. Compounds Containing Four Nitrogen and One Sulfur Atom in the Molecule ......... 1351 1. Triazolothiadiazines.............................................................................................. 1351 References .................................................................................................................................... 1351

I. INTRODUCTION The present chapter concerns the electrochemistry of heterocyclic compounds and their benzo derivatives. It is divided into three main sections: (1) six-membered heterocyclic systems, (2) fivemembered heterocyclic systems, and (3) fused heterocyclic systems. Each one of these sections is divided into two subsections: electrosynthesis and electrochemical reactivity. However, this chapter will be mainly focused on aromatic heterocycles and some of their dihydroderivatives, avoiding completely saturated rings such as pyrrolidines, piperidines, pyrans, or tetrahydrofurans, which have been already treated in other chapters as the corresponding aliphatic derivatives. On the other hand, we present a revision on around the last 30 years; older literature is available in previous editions of this book. A number of reviews on different aspects of the electrochemistry of heterocyclic compounds are disposable in the literature [1].

II.

ELECTROSyNTHESIS OF SIX-MEMbERED HETEROCyCLIC COMPOUNDS

A.

COMPOUNDS CONTAINING ONE NITROGEN ATOM IN THE RING

1. Pyridine and Derivatives Few examples can be found in literature concerning with the electrosynthesis of the pyridine ring. One of the first reactions was performed by bringing together an alkyne and a nitrile using a

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1312

Organic Electrochemistry

cyclopentadienylcobalt complex catalyst, generated in situ during the reaction. In this case, acetonitrile reacted with acetylene to give picoline in 83% yield [2].

2 HC

Co complex Bu4NBF4 –2.7 V

CH + CH3CN

N

CH3

The formation of 2,6-dimethyl-4-arylpyridine-3,5-dicarbonitrile, as the major product, from a wide range of aromatic precursors in acetonitrile has been described [3]. It is assumed that the generated 3-aminocrotonitrile anion initially passes through a glass frit diaphragm to the anodic compartment. cathodic CH3CN

–

reduction

CH2CN

H3C

CH3CN

–

H3C

CN

CN NH2

NH i

–N

–

H3C

CN

Two concurrent pathways are proposed to explain the formation of the final product: Ph +

NC SCN

–2e– –H+

i

SCN

C

i –SCN–

NH2 Ph SCN

NC Ph NC

NC

i –SCN –

CN

Ph

i

+

C

–2e– –H+

C

NH2 NH2 –NH3

NH2

NH2

Ph

Ph NC

NC

CN

CN

–2e– –2H+

N H

N

An alternative paired mechanism has been postulated for the oxidation of toluene to the corresponding 2,6-dimethyl-4-arylpyridine-3,5-dicarbonitriles [3]. Another interesting example in the electrochemical synthesis of the pyridine ring has been described by paired electrolysis of only acetonitrile as substrate [4]. 2,4,6-Trimethylpyridine-3,5dicarbonitrile was the main product. 2,6-Dimethylpyridine-3,5-dicarbonitrile and 2,6-dimethylpyridine-3,4,5-tricarbonitrile were obtained in minor quantities. Me

CH3–CN

Electrolysis Bu4NHSO4

NC

NC

CN

+ N

CN

H NC

CN

CN

+ N

N

Product: 4.7 g/0.4 F (84% of product) (13% of product) (3% of product)

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1313

Heterocyclic Compounds

2. Quinoline and Derivatives When CHCl3 is reduced at a Pb cathode, it undergoes a decomposition cascade reaction to form dichlorocarbene: +2e

CHCl3

+CHCl3

CHCl2–

–Cl–

–Cl–

CCl3–

–CH2Cl2

: CCl2

Thus, the electrolysis of CHCl3 in the presence of 2,3-dimethylindole affords 3-chloro-2,4dimethylquinoline in 46% yield [5]. CH3

CH3

Cl

: CCl2 CH3 N H

CH3

–

N

Similarly, the reaction of 1,2,3-trimethylindole with electrogenerated dichlorocarbene yields 3-chloro-1,4-dimethyl-2-quinolone [5]. N-Oxyquinolines are prepared from o-nitrobenzoyl derivatives by cathodic reduction via cyclization of the intermediate hydroxylamine [6]. OH CO–CH2X NO2

+4e +4H+ –H2O

N

X = CN, COPh Y = NH2, Ph

Y

O

Cathodic reduction under aprotic conditions of N-(2-acyl(or aroyl)-phenyl)-2,2,2-trichloro-Nalkyl-acetamide at −1.2 V (vs. SCE) yields 3-chloro-1,4-disubstituted-2(1H)-quinolinones as the major product [7].

R O N

R O

+1e– N

CCl3

R O

Cl + N

CCl3

R

R Cl

–OH–

O

R΄ (70–82%)

OH – N R΄

R Cl O

+2e– – –Cl

– CCl2

R’

R΄

R,R΄= alkyl R΄

N

O

O

O

OH

N R΄

O—

R Cl Cl O

H+ N

Cl Cl O

R΄

Direct electrolysis of pure nitroalkane under an inert atmosphere was used for the synthesis of 4-alkylquinolines by addition of a β-(2-aminophenyl)-α,β-ynone to the cathode compartment at the end of the electrolysis, that was carried out under galvanostatic control in a divided cell equipped with platinum electrodes [8].

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1314

Organic Electrochemistry R2 R5 R4

Electrolysis

O

R4

R5

R1 R2

R3 NH2

NO2

NO2 R1

N R3

Anodic oxidation of benzoylacetanilide in CH3CN/LiClO4 solution using controlled potential conditions (+1.8 V vs. Ag/Ag+) gives 1-phenylquinoline-2,4(1H,3H)-dione [9]. O

O

O –2e–/–2H+ NH–Ph

O

N Ph

3. Isoquinoline and Derivatives The electrochemical reduction of iminium salts in the presence of bromobenzyl bromides gives 1-(bromobenzyl)isoquinoline derivatives in moderate yields. This reaction is useful in the synthesis of several alkaloids [10]. Salsolinol-1-carboxylic acid (A) was electrochemically oxidized to the corresponding orthoquinone intermediate 1,2,3,4-tetrahydro-1-methyl-1-carboxy-6,7-isoquinoline dione (B) in pH ═ 7 aqueous solution. The ortho-quinone B is extremely unstable and rapidly decarboxylated to generate C, that can be further oxidized to give 1-methyl-6,7-isoquinolinediol [11].

HO

–2e–/–2H+

O

CO2

NH

Me

Me COOH B

HO

N

O Me

+

–2e /–1H +

O

–H+

N HO

Me

C –

O

HO

HO

HO

O Me COOH A

N

NH

NH

HO

HO

O

NH

N

Me

Me

Me

+

O

HO

Anodic oxidation of 3,4-dimethoxyallylbenzene and iodine in CH3CN containing LiClO4 yields 1,3-dimethyl-6,7-dimethoxyisoquinoline. It is suggested that CH3C+ ═NI is an intermediate in this reaction and gives aromatic electrophilic substitution prior to oxidative ring closure [12]. MeO MeO

+ CH3C — —N I

MeO

–H+

MeO

CH3 N CH3

N-Benzyl-enaminones are cyclized at the anode to isoquinolines [13]. O

O Undivided cell Anode, cathode/graphite

O O

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NH

NaClO4, MeOH

O O

N

1315

Heterocyclic Compounds

4. Phenanthridine and Derivatives Cathodic reduction of 2-(2′-chlorophenyl)-1-phenyl-substituted five-membered nitrogen heterocycles leads to the synthesis of phenanthridines [14]. DMF Pr4NBF4

N –

Ph

N

cathode/Hg pool

C1

B.

COMPOUNDS CONTAINING TWO NITROGEN ATOMS IN THE RING

1. Pyrimidines Electrolysis of α-chloro-ethylbenzene in acetonitrile/Bu4NBF4 at +2.9 V (vs. SCE), using a divided cell, afforded 2,6-dimethyl-4-phenylpyrimidine in 51% yield [15]. Cl

Cl

Cl –e–

–H+

+

Cl –e–

+

CH3CN Cl –2H N N

+

+

–Cl–

N N

Cl CH3CN

N +

2. Phenazines Anodic oxidation of 1-aryl-3-methyl-3-(4-methoxyphenyl)triazenes in acetonitrile/0.1 M Bu4NPF6 on Pt leads to the 2,7-dimethoxy-5,10-dimethyl-5,10-dihydrophenazine radical cation [16]. +

Me N

OMe

OMe

Anodic —N N 2Ar N— Me

oxidation

+ 2ArN2+ MeO

N Me

3. Quinazolines Controlled potential electrolysis of o-nitrobenzoyl acetamide at a mercury cathode leads to 4-hydroxy2-methylquinazoline 1-oxide, resulting from a ring closure of the phenylhydroxylamine intermediate. OH CO–NH–CO–Me NO2

+4e–, +4H+ –H2O Hg

CO–NH–CO–Me NHOH

–H2O

N N

Me

O

Similarly, other quinazoline derivatives are obtained by the cathodic reduction of various N-(2-nitrobenzoyl)amides [17].

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1316

Organic Electrochemistry

Electroreduction of 2-nitrobenzonitrile in the presence of a variety of alcohols leads to 2-alkylquinazolin-4(3H)-ones, via the incorporation of the alkyl chain of the alcohol during the cyclization [18]. O CN

R–CH2OH/H2SO4 Pt –0.7 V

NO2

NH R

N

Anodic oxidation of per-aryl ketenimines at a Pt anode leads, in a one-pot procedure, to quinazolines as a result of dimerization of the corresponding radical cations [19,20]. Ph

X X

Ph

2

Ph

N

Ph N

0.9–1.1 V (vs Ag/AgCl)

N

+

N

N Ph

Ph Ph

C.

Ph

Ph

X

Ph

X

X

COMPOUNDS CONTAINING THREE NITROGEN ATOMS IN THE RING

1. Triazines and benzotriazines 3-Phenylbenzo-1,2,4-triazine and other benzotriazines can be obtained by reductive ring closure of the appropriate o-(nitrophenylazo)phenylnitromethanes in an acetate buffer [21,22]. NO2

NH–N NO2

N N

Ph

Ph

N H

NH2

NH–N

+2e–, +2H+ –H2O –0.4 V

Ph

NO2

H N

O2

N

NOH

NH–N

+4e–, +4H+ –H2O –0.3 V

NO2

Ph

+6e–,+6H+ NH2

NH–N

–NH3

N

NH2

Ph

–0.8 V Ph

–2 H2O

3,4-Dihydro-1,2,3-benzotriazines have been obtained in only one electrochemical step using a cell in which there are two closely consecutive flow-through porous electrodes of opposite polarities [23]. 1,2,4-Benzotriazines are also obtained by the electroreduction of o-nitrophenyl hydrazides at a mercury cathode [24]. O

NH–N NO2

R

+4e–, +4H+ –H2O

O

NH–N

–2H2O R

N

NHOH +2e–, +2H+ –H2O

N R

–2e–, –2H+ O

NH–N R NH2

N

+H+ –H2O

H N NH N

R

N-Aminopyrazoles are oxidized by electrolysis in MeCN to give 1,2,3-triazines. The best yield in triazine is obtained when pyridine or water is added to the solvent [25].

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1317

Heterocyclic Compounds

D.

COMPOUNDS CONTAINING ONE OXYGEN ATOM IN THE RING

1. Pyrones and Derivatives Cathodic reduction of 2-chloro-2-phenylacetyl chloride in CH2Cl2/Et4NCl at a Hg cathode leads to the formation of phenylketene, and thence either a mixture of α- and γ-6-benzyl-3,5-diphenylhydroxypyranones through trimerization, or the phenylacetate derivative of the γ-pyrone through tetramerization [26]. Cl

+2e–

Cl

Ph

–

–Cl–

Cl

Ph

O Ph

O

O HO H Ph C

O

Ph

Ph

Ph

Ph

+

O C

Ph

O

Ph Ph

O

O

+ Ph

O

O Ph

O

OH

CO

Ph

Preparative electrooxidation reaction of tetra-arylcyclopentadienones at platinum in acetonitrile Bu4NPF6 yielded the corresponding α-pyrones as major product from an oxygen insertion reaction [27]. 2. Chromenes 4H-Chromenes were prepared employing an electrocatalytic chain method by the coelectrolysis of salicylaldehydes and CH acids [28], using an undivided cell in ethanolic solution with NaBr as electrolyte. At the cathode: 2ROH + 2e–

2R–O– + H2

In the solution: CH2(CN)X + RO–

–CH(CN)X

O–

O R1

+ ROH

H –CH(CN)X

R1

X CN

OH

OH R2

X R1

–OH–

X

CN

–CH(CN)X

NH2

O

CH2(CN)X

O

NH–

R2

R2

R2

X

R1

X

R1 +

N–

CH2(CN)X X

X O

O R2

R2

CN

R1

OH–

OH

R2

X

R1

CN

NH

+ –CH(CN)X

Similarly, the electrochemically induced catalytic multicomponent transformation of cyclic 1,3-diketones, aldehydes, and malononitrile in alcoholic solvents results in the formation of substituted 5,6,7,8-tetrahydro-4H-chromenes in 85–95% yields [29], employing resorcinol, malononitrile, and various aldehydes in propanol to produce 2-amino-3-cyano-7-hydroxy-4H-chromenes [30]. 3. Chromones The indirect oxidation of 2-hydroxychalcones to flavonoids uses tris(4-bromophenyl)amine in MeOH–CH2Cl2 as a mediator [31]. O

O –e– OH

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+ Ph Med

Med

O

Ph

1318

Organic Electrochemistry

The anodic oxidation of enaminones, prepared from 2-hydroxyacetylarenes and N,Ndimethylformamide dimethylacetal, gives bichromones in good yields [32]. O

O

2 R

OH

CH3CN/TEAP Pt gauze

Me

N Me

O + 2Me2NH

R

O

R

O

4. Coumarins and Derivatives [2]-Benzopyran[4,3-c][2]benzopyrano-6,12-dione is formed as side product in the cathodic reduction of 3,3-dichlorophthalide. Initially, elimination of chloride ion from the electrogenerated anion provides a carbene, which then rearranges to form an o-bisketene that finally dimerizes to the fused bicoumarin [33]. O

O

O

O

–Cl–

+2e –Cl–

O

O

– Cl

Cl Cl

O O

O

C C

O

O O

Efficient syntheses of 3-chloro-4-substituted coumarins have been achieved by the cathodic reduction of 2-acyl(or aroyl)phenyl trichloroacetates in MeCN/LiClO4 at a mercury cathode [34,35]. The postulated electrochemical mechanism involves two reduction steps [36].

R O R΄

O

–

O R

O

O

CCl3

+1e– –Cl–, –0.5 V

R –

R΄

CCl2

O

R΄

O

O

Cl Cl O

R +2e–, +H+ –0.8 V, –Cl–

Cl R΄

O

O

Electrochemical formation of 6H-dibenzo[b,d]pyran-6-one and 2-benzopyran-1(1H)-one has been performed by cathodic reduction of 9,10-phenanthrenequinone and 1,2-naphthoquinone in a divided cell in the presence of oxygen [36a]. 5. Isocoumarins 1-(2-Haloaryl)-1-methylepoxides are electrolyzed with CO2, in the presence of Ni(bipy)32+ · 2BF4− in DMF, employing an Mg, or an Al/stainless steel couple electrode, to give 3,4-dihydro-4-hydroxy4-methylisocoumarin [37]. O

Me

Me

OH

+e –X–

O

X O

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1319

Heterocyclic Compounds

E.

COMPOUNDS CONTAINING ONE NITROGEN AND ONE OXYGEN ATOM IN THE RING

1. Oxazines and Derivatives The electrochemical synthesis of 2-phenyltetrahydro-1,2-oxazin-3(2H)-one has been performed by electroreduction of nitrosobenzene in an aprotic solvent and in the presence of an equivalent of 4-chlorobutanoyl chloride [38]. Cathodic reduction of O2 in MeCN/TEAP in the presence of CO2 gives a carboxylating reagent (O2−•/CO2) that is able to convert amines and their derivatives into carbamates. In this way, ω-bromopropyl amine gave tetrahydro-1,3-oxazin-2-one in moderate yield [39]. 1,4-Benzoxazin-3(2H)-one is formed when o-nitrophenoxyacetic acid and some of its derivatives (X ═ OMe or NH2) are reduced in a protic medium at a controlled potential [40]. O

O O

X

NO2

O

O

+4e–, +H+ –H2O

X

–HX N OH

NHOH

X = OH, OMe, NH2

O

+2e–, +2H+

–H2O

O O

O –HX

X

N H

NH2

O

F. COMPOUNDS CONTAINING ONE/TWO SULFUR ATOMS IN THE RING 1. Thiopyran, Thiochromone, and Derivatives Cathodic reduction of phenacyl dithiopivaloates yields 2-tert-butyl-4H-1-benzothiopyran-4-one. A bromide ion acts as the leaving group [41]. O

O tBu

S

+e

S Br

tBu

S

G. COMPOUNDS CONTAINING ONE NITROGEN AND ONE/TWO SULFUR ATOMS IN THE RING 1. benzothiazines Cathodic cyclization of methyl 2-(2-nitrophenylthio)acetate to give 4-hydroxy-1,4-benzothiazin3(4H)-one is used as a key step in the syntheses of several naturally occurring hemiacetals [42]. COOMe S

+4e +4H+ Hg cathode

S

NO2

–0.4 V vs SCE H2SO4/MeOH

O N OH

2. Dithiazines α-Methylene benzylnitriles, bearing potential leaving groups attached to the terminal position of the double bond, undergo reduction at a carbon–sulfur electrode to give different dithiazines depending on the nature of the solvent employed [43].

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1320

Organic Electrochemistry Ph

Ph

H

NC

X

+e –1.9 V DMF

S

O N

N

S

Me

H

Ph

X = Cl, OTs, OCOPh

Ph

O S

+e –1.9 V NMP (N-methylpyrrolidone)

N

N

S

+

S

O

N

Me

N

S

Me

III.

ELECTROCHEMISTRy OF SIX-MEMbERED HETEROCyCLIC COMPOUNDS

A.

COMPOUNDS CONTAINING ONE NITROGEN ATOM IN THE RING

1. Pyridine and Derivatives The synthesis of unsymmetrical biaryls from pyridyl halides and 2,6-di-tbutylphenoxide can be effected by an electrochemically induced SRN1 reaction in liquid ammonia [44,45]. O– Cl

Mediator 2,2΄-dipyridyl

+

OH

Liquid ammonia (–38°C)

N

N 84% yield

Electroreduction of mixtures of 2-pyridyl halides and aryl halides (substituted by an electronwithdrawing group) using a sacrificial iron anode and a nickel 2,2′-bipyridine complex catalyst in DMF as solvent led to 2-arylpyridines [46]. X

+e + N

EWG

Y

DMF

N

EWG

Similarly, pyridinecarboxamides can be cathodically reduced to amines, or alcohols, depending on the reaction conditions. The reactions are carried out industrially [47]. H2O–H2SO4 N

CONH2

CONH2

Pb cathode Yield: 86% Conversion: 95%

N

CH2–OH

CH2NH2

H2O–H2SO4(Ti3+ salts) Pb cathode Yield: 75%

N

N

Significantly, the electrochemical reduction of heterocyclic nitriles can compete with catalytic methods, and for example, aminomethylpyridines can be obtained by electrochemical reduction of the corresponding pyridyl nitriles [48]. CN

N

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CH2NH2

H2O–H2SO4 (Fe2(SO4)3) Pb cathode Yield: 55%

N

1321

Heterocyclic Compounds

Cathodic reduction of 2- and 4-vinylpyridine gives trans-1,2-di(heteroaryl)cyclobutanes as major product. They arise via radical anion-substrate cycloaddition [49]. It has been shown that the removal of the alkylsulfonyl protective group from 4-(1-adamantyl)2-(alkylsulfonyl)-pyridine or 2-(1-adamantyl)-4-(tert-butylsulfonyl)-pyridine can be carried out quantitatively by an electrochemical method. Therefore, the use of this group as the protective and activating one in substitution reactions to the pyridine ring seems very perspective [50]. SO2Alk +2e– +2H+ Ad

Ad

N

N

+ HSO2Alk

Electrochemical reduction of pyridine-3-carboxylic acid has been studied at amalgamated electrodes under acidic conditions using potentiostatic and galvanostatic techniques to produce pyridine-3-aldehyde [51]. The electrochemical oxidation of 2-methylpyridine to picolinic acid has been carried out in a divided cell on the industrial scale [52]. Indeed, the optimization of this process has been carried out using factorial-designed experiments [53]. H2O–H2SO4 CH3

N

Current efficiency: 67%

PbO2 anode Yield: 80%

COOH

N

A similar process for the oxidation of 4-picoline uses electrogenerated Co(III) acetate as the oxidant, but the product is 4-(diacetoxymethyl)pyridine [54]. Me

N

C anode, 4e–/molecule Co(OAc)2 (10 mol%) Ac2O, KOAc, 80°C

(AcO)2HC

N

4-Fluoropyridine was synthesized from pyridine at a platinum anode and potentiostatic conditions in CH3CN solution containing Et3N · 3HF as electrolyte and fluorine source [55]. Anodic oxidation of hydrazones in the presence of pyridine led to the synthesis of s-triazolo[4,3-a] pyridinium salts [56]. Ar’ NH Ar

N

Ar΄

N

–4e– –3H+

+

N N –

N

Ar

Electrochemical coupling of mono- and dihalopyridines catalyzed by nickel complex in undivided cell affords unsymmetrical 2,2′-bipyridines in moderate to good yield [57]. [Ni(bpy)]Br2, e–

R N

Br

TBABF4, DMF Zn anode

R

R N

N 25–75%

2. Quinolines and Isoquinolines Chlorinated quinolines and isoquinolines may be carboxylated without loss of chlorine by reduction in aprotic medium in the presence of CO2 [58]. The anion of 1-isoquinolinecarboxylic acid can be decarboxylated under reductive electrochemical conditions in methanol solution. The yield of the decarboxylation process increases in the presence of CO2. This yield also depends on the pH of the solution, decreasing when the alkalinity increases [59].

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1322

Organic Electrochemistry

A cathodic electrochemical regiospecific hydroxylation of isoquinoline and quinoline, via their carboxylic acid, has been described. The reaction is carried out at a graphite cathode and in the presence of water. The most important feature of this process is that only one hydroxylated product is obtained from each carboxylated derivative [60]. The main product in the electrocatalytic hydrogenation of quinoline and quinoline N-oxide was 1,2,3,4-tetrahydroquinoline, whereas that in the electrochemical hydrogenation of 8-oxyquinoline, and its dibromo derivative, was 1,2,3,4-tetrahydro-8-oxyquinoline [60a]. The electrochemical cyanation of N-alkyl-1,2,3,4-tetrahydroquinolines occurs regioselectively at the heterocycle, and the position of substitution depends on the nature of the N-alkyl substituent. The reaction is carried out in a flow cell fitted with a graphite felt anode [61].

B.

COMPOUNDS CONTAINING TWO NITROGEN ATOMS IN THE RING

1. Pyridazine and Derivatives In DMF 3,6-diphenylpyridazine may be reductively carboxylated [58]. The two-electron reduction of pyridazines, substituted by electron-withdrawing groups, leads to their corresponding 1,2-dihydroderivatives. Depending on the nature of the ring substitution, these intermediates can either rearrange into 1,4-dihydropyridazine isomers or be further electrochemically reduced, and ring contracted, to pyrroles [62]. COOMe

COOMe R

N N

R΄

COOMe

COOMe +2e, +2H+ Controlled potential

R

R

N N H

N H R΄

N H

R΄

COOMe

COOMe +2e, +2H+

R΄

R

MeOOC

COOMe

N H

A cross-coupling reaction between 3-chloro-6-methoxypyridazine and methyl 4-bromobenzoate took place in DMF under constant current conditions with an iron rod as the anode and 10% of NiBr2bpy as the catalyst. Methyl 4-(6-methoxypyridazin-3-yl)-benzoate was obtained in 57% yield by employing a stoichiometric amount of both reagents [63].

MeOOC

+e–, NiBr2bpy Fe anode MeOOC OMe DMF, 0.2 A

Br + Cl N

OMe N

N

N

1,2-Dihydropyridazine-3,6-dione was oxidized to pyridazine-3,6-dione that is unstable and converts to maleic acid [64]. HN O

OH

NH +2 H2O O

–4e–/–4H+, –N2

O

HO O

3-Amino-6-arylpyridazines can be obtained in good yield by nickel-catalyzed electrochemical arylation of 3-amino-6-chloropyridazines using an iron/nickel electrode as the sacrificial anode at room temperature [64a].

© 2016 by Taylor & Francis Group, LLC

1323

Heterocyclic Compounds

2. Pyrazine In aprotic electrolytes pyrazines and other diazaheterocycles can be reductively carboxylated at the nitrogen atoms [65]. COOH N

N

CO2–MeCN/Et4NBr Fe cathode

Selectivity: 78% Current efficiency: 75%

N COOH

N

Chloropyrazine adds one electron at a mercury pool cathode, in CH3CN containing 0.1 M TEAP, to yield a short-lived radical anion, from which chloride ion is expelled to give the pyrazinyl radical that can acquire a hydrogen atom from the solvent to afford pyrazine [66]. N

Cl

N

–Cl–

+H N

N

N

N

N

N

Cl –

+e–

3. Pyrimidine and Derivatives 2-Arylpyrimidines have been obtained from 2-chloropyrimidine and various functionalized aryl halides by electroreduction using an iron rod as the anode and a catalytic amount of nickel-bipyridine complex in a mixture of DMF and pyridine as solvent [46]. N X

N

+ N

FG

Cl

+e– DMF

N FG

FG: Electron-donating or electron-withdrawing group

Electrolyses of silylated pyrimidines, with protected arylthioribose, in the presence of a catalytic amount of NBS or Br2 as mediator, in an undivided cell afford the corresponding N-glycosides (nucleosides) in good yield [67]. 2-Pyrimidinethiol, 4-methyl-2-pyrimidinethiol, and 4,6-dimethyl-2-pyrimidinethiol afford by anodic oxidation in acetonitrile the respective disulfide in excellent yields [68]. 2,2′-Dipyrimidyltrisulfide may be prepared after the addition of 2-chloropyrimidine to a solution containing the S42− anion in acetonitrile (from a sacrificial sulfur cathode), and then anodic oxidation of the intermediate thiolate anion [69]. 4. Cinnolines Cyanocinnolines can be selectively converted into cinnol-4(1H)-ones by electrochemical reduction with a Pt electrode [70]. –

NC OO

CN R N

N

© 2016 by Taylor & Francis Group, LLC

+2e– PhCN, n-Bu4NBF4 O2, H2O

O R

N –

N

R –HN — —C — —O N H

N

1324

Organic Electrochemistry

5. Phthalazines and Derivatives In aprotic media, the electron transfer to the molecule of 1,4-dichlorophthalazine involves the anionic elimination of chloride ions, the cleavage of the pyridazine ring and the formation of phthalonitrile [71]. Cl N +2e– –2Cl–

N N

N

Cl

6. Quinoxalines The reduction of 2,3-dichloroquinoxaline proceeds as a successive elimination of two chlorine atoms, yielding a nonsubstituted quinoxaline [72]. N N

Cl

N

Cl

+2e–

–Cl– Cl E1/2 = –0.70 V

N

N +2e– –Cl–

N

E1/2 = –0.87 V

7. Phenanthroline and Derivatives 1,10-Phenanthroline 5,6-dione was transformed into the corresponding 1,3-dioxol in a single-step process by cathodic reduction in dichloromethane [73]. The anodic oxidation of 1,10-phenanthroline to its 5,6-quinone was performed over a platinum anode at constant potential of +2.4 V (vs. SCE) with a consumption of ca. nine electrons/substrate molecule, using an MeCN +0.8 v% water solvent, with 0.1 M NaClO4 [74].

C.

COMPOUNDS CONTAINING THREE NITROGEN ATOMS IN THE RING

1. Triazines and Derivatives 1,2,4-Triazines are reduced in DMF to their dihydroderivatives [75]. This even occurs when the triazine ring is embedded in a polycyclic array. For example, the electrochemical reduction of benzopyrano[1]pyrimido-[1,2-b][as]triazinones leads to the 5,12-dihydroderivatives [76]. H

O N

O R2

N

O

N

N

N

R1

R2

N

O

+2e +2H+

N

N

R

R1

H

R

Electrochemical reduction of 3,5,6-triphenyl-1,2,4-triazine, when carried out at pH ═ 3.6, furnishes 3,5,6-triphenyl-2,5-dihydro-1,2,4-triazine and 3,4,5-triphenylpyrazole [77]. Ph

N

NH

–0.7 V

+ Ph

Ph

N

Ph

Ph

N H

© 2016 by Taylor & Francis Group, LLC

Ph

Ph

N

Ph N

Ph

N N H

1325

Heterocyclic Compounds

An efficient and convenient electrosynthesis of (3-methyl-4,6-dioxo-4H-thiazolo[2,3-c][1,2,4] triazin-7-ylidene)acetic acid methyl ester was obtained by oxidation of 3,4-dihydro-6-methyl-3thioxo-1,2,4-triazin-5(2H)-one with acetylenedicarboxylic acid methyl ester at a constant controlled potential [78]. O

O

H N

S

O

–/–1H+

N

–1e NH

N

COOR

ROOC

D.

COOR

S

N

N

COMPOUNDS CONTAINING ONE OXYGEN ATOM IN THE RING

1. Pyrylium Salts 2,2′,6,6′-Tetraaryl-4,4′-bipyranylidenes have been synthesized by electrochemical reduction of the corresponding 2,6-diarylpyrylium salts. The electrosyntheses are conducted in a divided cell (cathode: gold gauze; anode: platinum gauze) [79]. Ar +

Cathodic reduction

O Ar

Ar

Ar

O

O

Ar

Ar

The reduction of pyrylium salts in the presence of an alkyl iodide yields 4-alkylated-4H-pyran, together with the corresponding dimer [80]. 2. Pyrones Electrochemical reduction of 4-pyrones, substituted at positions C-2- and C-6, in MeCN–H2O 80:20, leads to complex “double dimers” with a degree of stereoselectivity [81]. 3. Chromones When the anodic fluorination of 2-phenylflavone is carried out with a platinum electrode in anhydrous acetonitrile and Et3N · 3HF, as supporting electrolyte, 3-fluoro-2-phenylflavone is formed preferentially [82]. O

O

O

F

–2e 30°C Ph Et3N·3HF

O

Ph

4. Coumarins Enantioselective reduction of the prochiral enone 4-methylcoumarin in the presence of alkaloids in methanol/buffer solution gave chiral 4-methyl-3,4-dihydrocoumarin and an optically inactive hydrodimer. O O

O

MeOH-buffer Alkaloid (+e)

O

O +

* Me H O

© 2016 by Taylor & Francis Group, LLC

O

O

1326

Organic Electrochemistry

By carrying out the electrolysis at lower pH and using yohimbine as catalyst, the optical yield was increased to 47%, but the chemical yield was only 18% [83]. More recently, optimization of a number of experimental parameters, such as pH, supporting electrolyte, working potential, and concentrations of substrate and yohimbine, has afforded (R)-4-methyl-3,4-dihydrocoumarin in an ee of 67% and 36% of chemical yield [84]. The electrochemical oxidation of catechols in the presence of 4-hydroxycoumarin led to substituted 8,9-dihydroxy-6H-benzofuro[3,2-c]chromen-6-ones [85].

OH + OH

OH

O

OH R

–4e–/–4H+ O

R

O

O

OH

O

Reduction of 4-(bromomethyl)-2-oxo-2H-chromen-7-yl acetate at carbon cathodes produces a dimer 4,4′-(ethane-1,2-diyl)bis(2-oxo-2H-chromene-7,4-diyl) diacetate in 72–80% yields [85a].

E.

COMPOUNDS CONTAINING ONE/TWO SULFUR ATOMS IN THE RING

1. Thiopyrans and Derivatives Electroreduction of tetra-activated 4H-thiopyrans leads selectively to a mixture of four 5,6-dihydro2H-thiopyran diastereomers. The relative percentages depend on the experimental conditions [86]. S

MeOOC

COOMe

+4e–/+4H+, –Me2NH

MeOOC

S

COOMe

Controlled potential MeOOC

COOMe

COOMe

MeOOC

NMe2

Anodic oxidation of 2,4-diphenyl-6H-cyclopenta[b]thiopyran affords the corresponding thiopyrylium cation [87]. Ph

Ph

Ph

+ –H+ –e–

–e– H

S

Ph

H

S

Ph

S

Ph

–

2. Dithiines and Derivatives Anodic oxidation of 2,5-diaryl-1,4-dithiins in the presence of 2,6-lutidine gives a 2,2′-dimer in low yield (8–22%) [88]. Ar S

Ar –2e

Ar

S

–

Ar S

S S S

–2H+ Ar

Ar

F. COMPOUNDS CONTAINING ONE NITROGEN AND ONE SULFUR ATOM IN THE RING 1. Thiazines and Derivatives The nitration of phenothiazine in acetonitrile in the presence of excess NaNO2 produces 3-nitrophenothiazine in 90% yield [89].

© 2016 by Taylor & Francis Group, LLC

1327

Heterocyclic Compounds

The electrochemical reduction of 5-acetyl-2-phenyl-6H-1,3-thiazine when carried out at a mercury cathode in acetate buffer/ethanol (1:1) affords either the hydrodimer (80%) or 5-acetyl-2phenyl-3,6-dihydro-2H-1,3-thiazine (75%) depending on the applied potential [90]. S

Ph

S

Ph

S

Ph

Ph

NH

HN

S

–0.8 V (vs. SCE)

–1.4 V (vs. SCE) N

HN

O

O

O

O

If C-4 carries an electron-withdrawing group and C-5 an electronically neutral group, cleavage of C6-S occurs with formation of a substituted benzothioamide [91]. O

S

Ph

O N

+2e +2H+

NH

Ph

COOEt O

S

O COOEt

Controlled potential electroreduction (protic medium, mercury cathode) of substituted 2-ethoxy and 2-phenyl-4H-1,3-thiazines leads to 6H-1,3-thiazines and (or) pyrroles. The nature of the isolated products appears to be strongly dependent upon the pH of the medium and the type of substitution [92]. MeOOC

COOMe

+4H+

(Ph)EtO

S

(Ph)EtO +4e–

COOMe

N

COOMe

R N H

S

(Ph)EtO +2e– +2H+

N

NMe2

R

COOMe

COOMe R

Anodic fluorination of 3-aryl-2H-1,4-benzothiazine derivatives in dimethoxyethane/Et3N·3HF led to the corresponding 2,2-difluorobenzothiazine. The reaction goes through the dimer 2,2′-bis[3-phenyl-2H-1,4-benzothiazine] [93]. N N

Ar

Ar –1e–/–1H+

N

Ar

S

F F

–6e–/–2H+

S S

S

Ar

N

IV. ELECTROSyNTHESIS OF FIVE-MEMbERED HETEROCyCLIC COMPOUNDS A.

COMPOUNDS CONTAINING ONE NITROGEN ATOM IN THE RING

1. Pyrrole and Derivatives The electrochemical two-electron reduction of pyridazines, substituted with electron-withdrawing groups, primarily leads to their corresponding 1,2-dihydro-derivatives, which can undergo electrochemical reduction to give rise to activated pyrroles by a ring contraction reaction with extrusion of ammonia [94].

R΄

R

R

R N

+2e–

N

+2H+

R΄

R

© 2016 by Taylor & Francis Group, LLC

NH NH R

+2e–

R΄

R΄ NH2 NH2

+2H+ R

R΄ –NH3

R –NH3

H2N

N H

R

R

N H

R

1328

Organic Electrochemistry

Similarly substituted pyrroles can be obtained by sulfur extrusion from 1,3-thiazine heterocycles by an electrochemical reduction process [95]. Following this methodology, the cathodic reduction of 3,6-dipyridinylpyridazines provided 3,6-dipyridinylpyrrole of biological interest [96]. An extension of this “green” methodology, is the efficient access to 2,6-bis[5-(pyridine-2-yl)pyrrol-2-yl]pyridines by electrochemical ring contraction of the corresponding 2,6-bis[6-(pyridine-2-yl)pyridazin-3-yl] pyridines [97]. Cathodic reduction of phenacyl bromides N-acyl hydrazones leads to dimeric 1,4-diaryl-1,4-butanedione di-N-acylhydrazones, which give the corresponding 1-N-acylamino-2,5diarylpyrroles in good yields [98]. R2

R2

O

O

NH

NH

R1

N

N

DMF

Hg, divided cell Br

DMF/LiCIO4

R2

R1

R1

O

N H

R1

Reflux

N

R1

N NH O R2

Electroreduction of 2-phenyl-4H-1,3-thiazines in 0.5 M H2SO4/EtOH leads to the corresponding pyrrole in 88% yield [99]. S

Ph

MeOOC

COOMe

COOMe

–0.7 V/SCE 4e–/4H+

N COOMe

Ph

N H

NMe2

2. Indoles In acidic solution, cinnolines are reduced to dihydrocinnolines, which are reduced with cleavage of the nitrogen–nitrogen bond to give indoles [100]. H N

N N

CH3

CH3

CH3

CH3

H N

NH2 NH –NH 3

N +2e–/+2H+

+2e–/+2H+

Arenediazonium salts are easily converted to aryl radicals under controlled potential electrolysis. The aryl radicals are useful precursors of indolines, provided that an appropriate leaving group is present to pre-empt intermolecular reaction of intermediate radicals [101]. BF4–

N2 +

Br Br

N

R΄ +e– R

Ms

R΄

R΄ –Br N Ms

R

N Ms

R

R’

+H+ N

R

Ms

A new application of the electrogenerated cyanomethyl anion that involves its use as an electrogenerated base is employed in the synthesis of functionalized indoles from alkynyl anilines. The reaction is carried at 0°C in a divided cell equipped with Pt electrodes under galvanostatic control [102].

© 2016 by Taylor & Francis Group, LLC

1329

Heterocyclic Compounds R1

R5

R5 R4

+e–

CH3CN (0.1 M TEABF4)

R3

NHR2

R1 N H

R4 R3

In the reduction of 1-(2-chloroethyl)-2-nitrobenzene and 1-(2-bromoethyl)-2-nitrobenzene at C cathodes in DMF/TMABF4, in the presence of either phenol or 2,4-pentanedione, as a proton source, the only significant product is 1H-indole [103]. A convenient electrochemical approach has been developed for the synthesis of indole derivatives from catechols and α-oxoheterocyclic ketene N,O-acetals [104].

OH

t-Bu

O

NH +

Carbon rods NaOAc buffer

O

OH

OH

N

+0.2 V vs. Ag wire O

OH O

56%

Ph

Electrooxidation of indolinyl alcohols under constant current conditions in MeOH with MeONa and KI resulted in the dehydrogenation to the corresponding indolyl alcohol [105]. –e– +

N

N n = 2 or 3

–e–/–H+ Base

N

N

(CH2)n-OH

(CH2)n-OH

(CH2)n-OH

–H+ Base

+

(CH2)n-OH

The anodic synthesis of 2,3,3-triphenyl-3H-indole (88% yield) was performed in acetonitrile/ LiClO4 in a divided H-cell, with a platinum anode and a graphite cathode employing N-methyl-N(1,2,2-triphenylvinyl)benzenamine as substrate in the presence of 2,6-lutidine [106]. Ph Ph

Ph

NMe

N

N

Ph

Ph

+

–H

Ph OH

OH Ph H

Ph +

–e–

–2H+, +H2O Ph

Ph

Ph

–2e–

O

Ph

Ph Ph

Ph

–H+

Ph –e–

+

+ N

H

N

Ph

N

Ph O

Ph OH

H

Electrochemical synthesis of derivatives of 1H-indol proceeds by reduction of substituted o-nitrostyrenes at carbon cathodes in DMF in the presence of a 10-fold molar excess of a proton donor [106a]. Ph +4e– R NO2

© 2016 by Taylor & Francis Group, LLC

C cathode DMF, TMABF4

Ph

R N H

1330

Organic Electrochemistry

3. Carbazoles Carbazoles were successfully synthesized by oxidative cyclization of the corresponding diaryl derivatives using an electrochemically generated hypervalent iodine oxidant. Electron-withdrawing nitro and donating methoxy groups at the para position of the acetamide group interfered with cyclization [107].

R2

R1

Anodic oxidation R1 in CF3CH2OH I

R2 N

NHAc

B.

Ac

COMPOUNDS CONTAINING TWO NITROGEN ATOMS IN THE RING

1. Imidazole and benzo Derivatives The electrochemical reduction of phenacylazide thiosemicarbazones in aprotic DMF/LiClO4 medium, at a mercury cathode in a divided cell under controlled potential, leads to N-substituted imidazolethiones in a one-pot reaction with good yields [108]. N–NH–CS–NH2

2

N

+2e–

N3

–N3–

Ar

N

Ar

N

+2e–, +2H+

Ar N

N

Ar

Ar N H

S

S

Electrocatalytic reduction of phenacyl azides (only 10% of the expected current for a theoretical 2e−/molecule process) leads to the formation of 2-aroyl-4-arylimidazoles in 70–80% yield [109].

2

N3

Ar

N

Cathodic reduction

O

Ar

DMF/LiClO4

CO–Ar N H

Different Schiff bases, for instance, the one prepared from o-phenylenediamine and the corresponding aromatic aldehyde, were anodically oxidized in acetonitrile–TEAP as electrolyte solution at platinum and using controlled potentials. After two-electron oxidative cyclodehydrogenation, several imidazole derivatives were prepared in yields ranging from 60 to 90% [110]. The indirect electrooxidation of ketones in ammoniacal methanol using iodide ion as a mediator afforded 2,2-dialkyl-2H-imidazoles via an oxidative cyclocoupling of ketimine intermediates formed from ketones and ammonia [111]. R1

MeOH/NH3/KI NH

2 R2

I

+

–2e–

I

–

R1 R2

N N H

R2 R1

R1

N

R2

N

–2e–, –2H+

R2 R1

Anode

Electrooxidation of N-(4-methoxyphenyl)-N′-aryl-2,2,2-trifluoroethanimidamides in an CH3CN–H2O–NaClO4 –(C)–(Pt) system affords p-benzoquinone imines, which are converted to 1-aryl-2-trifluoromethylbenzimidazoles by acid-catalyzed cyclization [112].

© 2016 by Taylor & Francis Group, LLC

1331

Heterocyclic Compounds

N

O

HO

N

–2e–

N

Lewis acid

CF3 N

CF3

N

CF3

N H

Ar

Ar

Ar

MeO

Of particular interest is the fact that electrooxidation of the nonfluorinated imidamides in a MeCN–NaClO4 –(C)–(Pt) system also provided the desired benzimidazoles in good yields [113]. 2. Pyrazole and benzo Derivatives Electrochemical oxidation of the sodium salts of tropone and 2-phenyltropone tosylhydrazones afforded 2-tosyl-2H-indazole and 1-tosyl-7-phenyl-1H-indazole, respectively. The reaction proceeds through the cyclization of the corresponding hydrazyl radicals generated by electrochemical oneelectron oxidation of the hydrazone anions [114]. R

–

N N–Ts

–e–

N N–Ts

R

R

R

R

Na+

N

–H

N

N

N–Ts

N–Ts

N–Ts

Electrolysis in two steps (reduction then oxidation) at a mercury electrode of ortho-nitrobenzylamines was carried out to perform the synthesis of the nitroso derivatives. The electrophilic character of the nitroso group was used to generate the N–N bond of 2-substituted indazole. Only benzylamines bearing an electron-donating group were successfully cyclizated [115]. R

R

N H

N H

''Redox'' electrolysis

–H2O

N R

Hydroalcoholic medium

NO2

N

NO

1-Benzyl-1,2-dihydro-3H-indazol-3-one has been produced by electrochemical reduction/cyclization of N-benzyl-N-nitrosoanthranilic acid at a mercury cathode [116]. The electroreduction of several fused [1,3,4]-thiadiazinium salts, on a mercury pool cathode in an aprotic medium, formed a mixture of the corresponding fused pyrazoles and ketimines. The heteroaryl ketimine results from the opening of the 1,3,4-thiadiazine ring, whereas the fused pyrazole (72–87% yield) resulted from the ring contraction of the 1,3,4-thiadiazine ring [117]. SH S

Ph 3

+N

Ph

Ar N

N

S

Ph

Ph –4e– Ar Ph

+ 2

N

N

N

Ar

Ph

Pyrazoles with alkyl, cycloalkyl, or aryl groups are prepared by electrochemical oxidation of the corresponding 2-pyrazolines in the presence of ion-forming halides [118]. Anodic fluorination of spiropyrazole-5,3′-chroman-4-ones and their thiochromanone analogs in dimethoxyethane containing Et4NF·4HF resulted in ring opening of spiroheterocycles, which led to the formation of 5-(2-fluorocarbonyl)phenoxymethyl-1,3,4-triphenyl-pyrazole or their thioether analogs that are further oxidized to 5-formyl-1,3,4-triphenylpyrazole [119].

© 2016 by Taylor & Francis Group, LLC

1332

Organic Electrochemistry

H

Ph

Ar

O

Ph

Ar

O

–2e–, –2H+ Et4NF·4HF X = O,S

N N

F

–2e– H2O X=S

N N

X

Ph

X

Ph

Ar

N OHC

N Ph

Ph

Oxidation of p-substituted phenylhydrazones of 2-oxo-phenylacetonitrile yields derivatives of 1-phenyl-3-cyano-1H-indazoles [120].

C.

COMPOUNDS CONTAINING THREE NITROGEN ATOMS IN THE RING

1. Triazole and benzo Derivatives Anodic oxidation of heterocyclic hydrazones, prepared from 2-hydrazinopyridine and the corresponding aldehyde, was typically performed in an acetonitrile solution containing TEAP with the addition of 60% perchloric acid using controlled potential electrolysis [121,122]. H N

N N

Ar

–2e–, –2H+

N N

N

Ar

Arylazo-substituted enamines are converted into 2-aryl-1,2,3(2H)-triazoles by galvanostatic anodic oxidation in acetonitrile in fair yields without the use of metal ions [123]. Ar

N

N

CN

N

CN

N

Me

–2e–, –2H+ Ar

N

Me

H2N

Benzaldehyde phenylhydrazones are oxidized to nitrile imines in the presence of a pyridine; the two-electron oxidation product undergoes a 1,3-dipolar cycloaddition with the heteroaromatic system. The resulting triazole is further oxidized to a [1,2,4]triazolo[4,3-a]pyridinium cation [124]. Ph

Ph H N

N

N

–2e–, –2H+

H C

–2e–, –H+ N

N N

N

– CIO4

Ph

N

N

Ph

Electrochemical oxidation of N-thioamidohydrazones was undertaken at Pt electrode in acetonitrile providing 1,2,4-triazol-3-thione in a satisfactory yield [125]. H

N N

Ar

NH S

© 2016 by Taylor & Francis Group, LLC

R

NH

–2e–, –2H+

Ar

NH S N R

1333

Heterocyclic Compounds

2-(3-Aryl-5-methyl-1H-[1,2,4]triazol-1-yl)-5-aryl-1,3,4-thiadiazoles were obtained in moderate yields by anodic oxidation in acetonitrile of 2-arylidene-1-(5-aryl-1,3,4-thiadiazol-2-yl)hydrazine under galvanostatic conditions and at a graphite anode [126]. Ar N

N

N

Ar

+

N

Anodic oxidation N

Ar

N N

CH3CN H

+

–H+

N

N

Ar

N

N

Ar

N

N S

N N

N

S

N

N Ar

N

S

Ar N

N

Ar

–2e–/–H+

N H

S

Ar

S

Ar

Me

Me

C

Me

N

In the process, the benzyl cation is followed by a Ritter reaction with the solvent to give the product after subsequent intramolecular cyclization.

D.

COMPOUNDS CONTAINING ONE NITROGEN AND ONE OXYGEN ATOM IN THE RING

1. Oxazoles Electroreduction of N-acylated imidates leads to a cyclic 3a,6a-dihydrooxazolo[5,4-d]oxazole in good yield [127]. O R2

N

R2

N

+1e–

R1

O

R2

O N

Dimerization

N

Cyclization R1

OR3

R1

OR3

O

R1

R2

2. Isoxazole and benzo Derivatives Electrochemical reduction of nitroethylenic ketones in hydroorganic medium at pH ═ 1 forms 3,4,5triphenylisoxazole. The process involves an intermediate α,β-nitrosoethylenic compound, which gives an 1,3-diketone oxime upon reduction. This last compound cyclizes into an isoxazole [128]. O

O

O Ph

Ph Ph

R

Ph +2e–/+2H+ –H2O

Ph

R

NO2

Ph

Ph

+2e–/+2H+

–H2O O

R = Ph Ph

NO

Ph

Ph

NOH

N

Preparative electroreduction of appropriately substituted aryl-nitrones in aqueous alcoholic media at a mercury cathode and in a flow cell fitted with a graphite felt cathode gives as major product 3-methoxycarbonylbenzoisoxazole [129]. COOMe N

COOMe Electroreduction

NO2 Ph

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N C-addition Ph NHOH

MeOOC NH–Ph O N H

COOMe

–NH2Ph

O N

1334

E.

Organic Electrochemistry

COMPOUNDS CONTAINING TWO NITROGEN AND ONE OXYGEN ATOM IN THE RING

1. Oxadiazole 5-Substituted-2-amino-1,3,4-oxadiazoles were synthesized directly from the semicarbazone of different aldehydes at a platinum electrode under controlled potential electrolysis in an undivided cell assembly in acetonitrile [130]. Anodic oxidation of N-amidoimidates was performed in acetonitrile/LiClO4 at Pt using controlled potential. Oxidation of the amide group followed by deprotonation and intramolecular cyclization reaction led to 1,2,4-oxadiazol-3-one derivatives [131]. O

O NHR2

N R1

–2e– –H+ +H2O

OR3

O

N R1

NR2

Intramolecular

OH

cyclization –R3OH

OR3

N N

R1

O

R2

The one-pot concomitant electrochemical reduction of phenanthrenequinones and arenediazonium salts led to the formation of 1,3,4-oxadiazol-2-(3H)-ones and dibenzo-[c,e]-azepines when N-methylformamide was used as the solvent [131a]. O

X

O +

X

N

Ar

+e– –0.5 V N H–CO–NHMe

X

X

X

X

O

O NHMe

N N–Ar

+ N N O

O

N Ar H

Me

X = CH, N Ar = C6H5 (a), 4-MeO–C6H4 (b), 2-MeS–C6H4 (c), 4-CI–C6H4 (d), 4-Br–C6H4 (e), 4-MeCO–C6H4 (f )

Anodic oxidation of aldehyde-N-aroylhydrazone at a platinum electrode in an undivided cell provides the corresponding 1,3,4-oxadiazoles at ambient condition [131b].

F. COMPOUNDS CONTAINING ONE NITROGEN AND ONE SULFUR ATOM IN THE RING 1. Thiazole, Isothiazole, and benzo Derivatives 3-Aryl-2-phenylsulfonyl propenenitrile, used as vinyl sulfones in acetonitrile, were studied using a reactive sulfur–graphite electrode. Electroreduction of these compounds gave 5-aryl isothiazoles bridged with two or three sulfur atoms at the 3-position [132]. NC

Ph

SO2Ph

1) +1e– (S/C electrode) 2) Mel reflux

Ph +

Ph

S

S

NC

SSMe

N

+ S–S–Me

Ph

N S–Me

The reduction of N-thioamidoimidates in aprotic media at a mercury electrode leads to a dimer that evolves to thiazolo[5,4-d]thiazole by intramolecular cyclization [133]. Several benzylideneamino thiophenols were electrochemically oxidized in methanol containing sodium acetate as the supporting electrolyte to afford the corresponding 2-arylbenzothiazoles.

© 2016 by Taylor & Francis Group, LLC

1335

Heterocyclic Compounds

The reaction proceeds via intramolecular cyclization involving the formation of a new bond between the benzylic carbon of the substrate and the sulfur of the thiol group [134]. SH

S –2e–

Ar

–2H+ N

N

Ar

G. COMPOUNDS CONTAINING ONE OXYGEN ATOM IN THE RING 1. Furan and benzo Derivatives Electrochemical reduction of phenacyl bromides in dry DMF at −1.0 V yields 2,4-diarylfurans in good yields [135,136].

+2e–

Ar

Br Ar

Ar

–Br–

Ar

O

O

O

O

O Br

–H2O Ar

–Br–

Ar 65–80%

Ar

However, when the carbonyl group is protected by semicarbazone formation, its cathodic reduction leads to the dimeric product in quantitative yield. After hydrolysis and dehydration, 2,5diarylfurans are obtained [137]. On the other hand, when α-bromopropiophenone is reduced in aprotic medium, 1,4-diphenyl-2,3-dimethyl-1,4-butanedione is formed. Dehydration of that produces 2,5-diphenyl-3,4-dimethylfuran [138]. When phenacylbromides are reduced in dry acetone/LiClO4, 4-aryl-2-methylfurans are obtained [139]. O

O Br

Ar

O

O

O

+2e–

+

Ar

– O

O

+

Me

Ar

–Br–

O Br

Ar

Ar

O –H2O

–Br– Me

O

Ar

Me

Cathodic reduction of 2-bromo-2-cyanoacetophenone afforded 5-amino-4-benzoyl-3-phenylfuran-2-carbonitrile in a one-pot reaction [140]. Effective synthesis of tetraethyl furan-2,3,4,5-tetracarboxylate was described from electrondeficient alkynes via the catalysis of electrogenerated base and Fe3+ ions [140a]. Propargyl ethers in combination with catalysis by Ni(cyclam)Br2 were electrolyzed in DMF in a single-compartment cell, fitted with a consumable magnesium anode and a carbon fiber cathode under constant current. Intramolecular cyclization involving triple bonds and rearrangement produced 3-methylbenzofurans [141]. X

O

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1) +e– Ni(cyclam)2+ (10%) 2) Hydrolysis

Me

O

(X = Br, I)

1336

Organic Electrochemistry

Anodic oxidation of enol carbonates yields benzofuran after O–CO bond cleavage to provide α-carbonyl cation and acyl radicals respectively. Cyclization of the α-carbonyl cation, [1,2]methyl shift and deprotonation finally leads to the benzofuran [142].

Mes

Mes

O

Mes

R1

–COR2

R2

R1

R1

–e–

CH3 H3C

+

–e–

R2 Mes

O

Mes

O

Mes

+

O

O

+

H3 C

O

Mes

O

O +

[1,2]–Me

R1

R1

–H+

Mes

R1

Mes

CH3

Mes

CH3

CH3

H. COMPOUNDS CONTAINING ONE SULFUR ATOM IN THE RING 1. Thiophene and benzo Derivatives The cathodic reduction of phenacyl thiocyanate at controlled potential of −1.2 V, using DMF/ LiClO4, afforded 2,4-diphenylthiophene, but in low yield [143]. Anodic initiated introduction of sulfur and thiol substituents into five-membered heterocycles such as furan, thiophene, and pyrrole has been performed. That method allows activation of H2S at room temperature up to a radical cation and to transform furan and pyrrole into thiophene and 2-thiophenethiol [144].

I. COMPOUNDS CONTAINING ONE SULFUR AND ONE OXYGEN ATOM IN THE RING 1. Oxathioles Electrochemical reduction of monothiobenzils in the presence of carbonimidoyl dichlorides provides 4,5-diaryl-2-arylimino-1,3-oxathioles in good yield [145]. Ph

O Ph

Ph

+2e, ArN — — CCI2

NAr

–2CI– Ph

S

O S

V. ELECTROCHEMISTRy OF FIVE-MEMbERED HETEROCyCLIC COMPOUNDS A.

COMPOUNDS CONTAINING ONE NITROGEN ATOM IN THE RING

1. Pyrrole and Derivatives Cathodic reduction of 2-(2′-chlorophenyl)-1-phenylpyrrole in DMF gave pyrrolo[1,2-f]phenanthridine along with 1,2-diphenylpyrrole [14].

Cathodic reduction N CI

© 2016 by Taylor & Francis Group, LLC

Ph

N

1337

Heterocyclic Compounds

The electrochemical oxidation of 4,5-diphenyl-2-(p-tolyl)pyrrole at +0.45 V in neutral unbuffered acetonitrile affords, almost quantitatively, 1,6a-dihydro-8-methyl-2,3,4,5-tetraphenyl-6a(p-tolyl)benzo[g]pyrrolo[3,2-e]indole [146]. Ph

Ph R

+0.45 V R

Ph N

Ph

R

Ph R

Ph

N H

Ph

Ph

N

Ph

N H

HN

R = p–Tolyl

CH3

Ph

Electrochemically generated tetraethylammonium peroxydicarbonate and tetraethyl ammonium carbonate react under very mild conditions with pyrroles affording, after addition of a suitable alkylating agent, the corresponding N-alkylated pyrroles in high yield [147]. The electrooxidation of several 1-arylpyrroles has been carried out in methanol containing sodium cyanide at a platinum anode in a divided cell. As a result of two-electron oxidation, the corresponding pyrrole cyanides were obtained in yields ranging from 76% to 85% [148]. CN

R

R

–2e–, –H+ N

N

NaCN/MeOH

Anodic fluorination of 2-cyano-1-methylpyrrole using Et3N · 2HF in an undivided cell provides the corresponding 5-fluoropyrrole and 2,5,5-trifluoro-1-methyl-3-pyrrolin-2-carbonitrile, while the use of Et3N · 3HF afforded selectively the latter product, which was readily hydrolyzed to isolable 5,5-difluoro-1-methyl-3-pyrroline-2-one. This is the first report of successful anodic fluorination of a pyrrole derivative [149]. Potentiostatic electrolysis (divided cell, anode compartment) of 1-methyl and 1-phenylpyrroles in the presence of sodium 4-nitro-pyrazolate in MeCN/MeOH gave the corresponding 2-(pyrazol1-yl)-pyrroles [149a]. 2. Indole and Derivatives Preparative electrolyses of 5-nitroindoles in acidic and basic methanol solutions gave 4-substituted 5-aminoindoles resulting from the regiospecific addition, to a diiminoquinone intermediate, of methanol or methanolate ion and/or of any other good nucleophile present in the electrolytic solution [150]. R

R

N H

R

HOHN

O 2N

HN

+4e–, +4H+ –H2O

N H

–H2O

N

Nu H2N +NuH

R

N H

Indoles protected at the N−1 position with the N,N-dimethylaminosulfonyl group were efficiently deprotected by electrolysis [151]. R3 R5

R3

R5 +2e–, +2H+ DMF

+ HSO2NMe2 N

N R7

© 2016 by Taylor & Francis Group, LLC

SO2NMe2

R7

H

1338

Organic Electrochemistry

The electroreduction of 1-indolealkanones in isopropanol produces five-, six-, and seven-membered trans-cyclized products stereospecifically using a Pb cathode and a divided cell containing Et4NOTs [152]. X

X

OH

H

Me

O

+2e–

Me

N

N

iPrOH

n

n

2,3-Diphenylindole was oxidized to the dimer in 90–95% yield, and a radical coupling mechanism was proposed [149]. N

Ph

Ph Ph

Ph

NH

The anodic amination of 1-hydroxy-2-phenylindole in dry acetonitrile at the Pt electrode proceeds efficiently but only using amines with oxidation potentials lower than that of the substrate [153]. Reaction of anodically generated thiocyanogen with 3-alkyl-substituted indoles results preferentially in isothiocyanation at the indole 2-position rather than in the expected thiocyanation [154]. The anodic oxidation of indoles and alcohols in the presence of cyclodextrins gave diindolylmethanes in good yield [155]. R3

R3 + R4–CH2OH N R2

R1

R3

R4

CD/–e– R2 NaCIO4/H2O R2

N R1

N R1

[4+2]-Cycloaddition reactions between 2-vinylindoles acting as hetero-dienes and β-acceptor substituted cyclic and acyclic enamines can be induced by formation of 2-vinylindole radical cation via anodic oxidation. Pyrido[1,2-a]indoles or indolo[1,2-a]hexahydro-1,8-naphthyridines are formed in one step with complete regio- and stereochemical control [156]. 3-Alkylindoles have been synthesized anodically by coelectrolysis of indole-3-propanoic acid and monocarboxylic acids Me–(CH2)n –COOH (n ═ 1, 4, 7, 8, 10, 14). The desired product is obtained in 33–60% yield at Pt anode [157]. Anodic fluorination of various N-acetyl-3-substituted indole derivatives was successfully carried out in Et4NF·4HF/MeCN to provide the corresponding trans-2,3-difluoro-2,3-dihydroindoles exclusively or selectively [158]. The electrochemical oxidation of the central mammalian alkaloid 1-methyl-6-hydroxy-1,2,3,4tetrahydro-β-carboline has been studied in neutral aqueous solution at a pyrolytic graphite electrode. Voltammograms of this alkaloid show two closely spaced oxidation peaks. Electrolysis at the less positive potential produces a radical intermediate that dimerizes to give two diastereomers of 5,5′-bi(1-methyl-6-hydroxy-1,2,3,4-tetrahydro-β-carboline). At potentials more positive the putative radical intermediate is further electrooxidized to a C(5)-centered carbocation, which reacts with the substrate in an ion-substrate reaction to afford the dimer or, with water, to give ultimately 1-methyl1,2,3,4-tetrahydro-β-carboline-5,6-dione [159].

© 2016 by Taylor & Francis Group, LLC

1339

Heterocyclic Compounds H NH

HO

–, –H+

–e

NH

O

Me

N H

–e–, –H+

+

O

NH

Me

N H

Me

N H

Low potential

High potential

N H

Me

H

O NH

O HN

OH NH

HO N H

Me

N H

Me

5-Hydroxytryptamine (5-HT, serotonin) is oxidized to a C(4)-centered carbocation intermediate, which reacts either with 5-HT, to give 5,5′-dihydroxy-4,4′-bitryptamine as the major product, or with water to give tryptamine-4,5-dione [160]. +

O

+

HO

NH3 –H+–2e– +H++2e–

N H

OH

+

NH3

+

HO

NH3

+H2O

N H H+

5-HT

N H +H+, +2e– –H+, –2e– +

OH

HN +

+

O

NH3

NH

HO

H3N

O

NH3

N H

The electrochemically driven oxidation of 5-hydroxytryptophan in the presence of free glutathione yields 4-S-glutathionyl-5-hydroxytryptophan and 7-S-glutathionyl-tryptophan-4,5-dione [161]. COOH HO

COOH

NH2 –H+ –2e–

O

NH2

COOH

OH HO

NH2

+H2O –H+

+H+ +2e– N H GHT

N H GSH –4H+ –4e–

N H H+ COOH

SG HO

NH2 N H

COOH

O O

NH2 N H SG

When the anion of indole-3-acetic acid, a growth hormone in plants, is electrochemically oxidized (one electron), an acetoxy radical is formed. This radical undergoes a second one-electron oxidation/decarboxylation to a carbocation precursor of 3-hydroxymethyl-2-oxindole, indole-3carbinol and 3-methylene-2-oxindole [162].

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1340

Organic Electrochemistry

3. Indolizines Oxidation of 2,3-diphenylindolizine-7-carbonitrile was performed at a Pt mesh electrode. Initial oxidation of the neutral indolizine gives rise to the corresponding indolizine radical cation that undergoes dimerization [163]. CN

CN

CN

CN Ar

Ar

Ar

Ar

H +

–2e–

N

+ N

2

2 N

Ar

B.

Ar

N

N

H Ar

Ar

–2H+

N+

Ar

Ar

Ar

Ar CN

CN

COMPOUNDS CONTAINING TWO/THREE NITROGEN ATOMS IN THE RING

1. Imidazole and Derivatives N-alkylation of imidazole was easily performed in 77% yield by the electroreduction of the substrate in the presence of EtOMs [164]. The electrooxidation of several 1-methylimidazoles was performed in methanol that contains NaCN at a Pt anode in a divided cell. The replacement of an aromatic hydrogen by a cyano group occurred at a position lacking the substituent in the imidazole ring. 1-Methylimidazole gives three possible ring-substitution products: 5-, 2-, and 4-cyano isomers in an approximate ratio of 5:2:1 [165]. Imidazole-2-(3H)-thiones were electrooxidized to the corresponding disulfides [166]. R3

NH

R2

N R1

R3 EtOH/2M HCl Pt-Pt S –2H+, –2e–

R1

N

N

S R2

N R1

R2

S N 54–90%

R3

Synthesis of imidazolium carboxylate compounds was efficiently achieved via carbenes by electrochemical reduction of imidazolium precursors under very mild conditions [166a,b]. Similarly these N-heterocyclic carbenes can be trapped by selenium to yield a selenourea derivative [166c]. 2. Pyrazole and Derivatives An electrochemically induced catalytic tandem Knoevenagel–Michael reaction of two equivalents of 5-methyl-2-phenyl-2,4-dihydro-3H-pyrazole-3-one with various aromatic aldehydes in ethanol in an undivided cell in the presence of sodium bromide as an electrolyte results in the formation of the corresponding 4,4′-(arylmethylene)bis(1H-pyrazol-5-ols) in 80–96% yield [167]. Me

O

N

R +

EtoH, NaBr

N Ph

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Electrolysis 0.04e–/molecule

CHO

R Me

HO Ph

N N N

N Me

HO

Ph

1341

Heterocyclic Compounds

Anodic oxidation of ethyl 5(3)-amino-3(5)-(phenylamino)pyrazole-4-carboxylate at a Pt electrode first gave the cation radical, which underwent deprotonation and dimerization [168]. EtOOC

H2N

NH–Ph

Ph–HN

COOEt

N

–e–, –H+

N

N

N

H2N

N H

N

EtOOC

NH–Ph

NH2

The electrooxidation of 1-methylpyrazole in methanol containing sodium cyanide produced 1-methylpyrazole-4-carbonitrile and 5-carbonitrile in yields of 23% and 8% respectively [169]. A bis-ketenedithioacetal disulfide was obtained by anodic oxidation of 1-phenyl-3-methyl-4butyldithiocarboxylate-5-pyrazolone using a glassy carbon electrode [170]. S

Me

Me

Ph N

S

0.73e–/molecule 2

O

SBu

SBu

S N

N

HN N

O

EtOH/H2O (85:15) LiCI

N

BuS O

Me

Ph

Ph

Electrolysis of a mixture of 3,5-dimethylpyrazole and 1,4-dimethoxybenzene in acetonitrile afforded ortho-substitution and ipso-bisaddition in 28% and 16% current yields respectively. It should be noted that N-arylation of pyrazole, performed under analogous conditions, gave rise only to the ipso-bisaddition product in 25% current yield [171]. Me

OMe

Me

2e–/molecule N

Me

+

Bu4NCIO4 OMe

Me N

OMe

Me

Me

N + N

N H

Me

OMe

N

N

N

Me MeO

OMe

4-Halogenated pyrazolecarboxylic acids were synthesized via halogenation of the corresponding acids at a Pt anode in aqueous solution of alkaline halides under conditions of divided galvanostatic electrolysis [172]. Electrochemical oxidation of di- and trisubstituted pyrazole-4-carbaldehydes on an Ni-anode in aqueous alkali led to the formation of the corresponding pyrazole-4-carboxylic acid in 60–90% yield [172a]. Similarly, oxidation of pyrazoles containing hydroxyethyl group bonded with the pyrazole ring at its N and C atoms afforded pyrazole-1-acetic acid or pyrazole-4-carboxylic acid respectively [172b]. 3. Triazoles and benzotriazoles N-arylazoles were prepared through a paired electrolysis starting from a mixture of benzene, 3-nitro-1,2,4-triazole (NTA) and an Me4N+ salt of NTA in an undivided cell [173], or starting from a mixture of 1,4-dimethoxybenzene with an Me 4N+ salt of NTA in the anodic space of a diaphragm cell [174].

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1342

Organic Electrochemistry ++++++ –e–

–e–, –H+ +

N

N

N

N

NO2

NO2

NO2

N

N

N

N H N

N

N

N

NO2 +e–, −1/2H2 ––––––––

1-Phenylbenzo-1,2,3-triazole is reduced in aprotic medium to give a mixture of carbazole and diphenylamine. Reduction of benzotriazole in the presence of an excess of acetic anhydride yielded 1,2,3-triacetyldihydrobenzotriazole [175].

C.

COMPOUNDS CONTAINING ONE OXYGEN/SULFUR ATOM IN THE RING

1. Furan and benzo Derivatives Electrolytic reduction of 2-(1-bromo-1-methylethyl)benzofurans in acetonitrile affords the corresponding 2,2-dimethylchromenes in good yields even in the absence of a proton donor and comprises the cleavage of a carbon–bromine bond followed by ring expansion [176].

Me

–2e– –Br

R

O

+H+ Me

–

Br

R

O

Me

Me

Oxidative ring opening of 2-substituted furans into methyl 4-substituted(E)-4,4-dimethoxy-2butenoates has been performed by electrolysis in an NH4Br–Et4NClO4 –MeOH–(Pt electrodes) system. The electrogenerated Br+ plays an important role in the ring-opening reaction [177]. MeOOC OMe –2e–, Br–

–2e– MeO

Y

MeO OMe

Y

Y

O

O

2,5-Dimethoxy-2,5-dihydrofuran was prepared by the oxidation of furan and by the reduction of methanol solvent, using a thin-layer flow cell geometry with anode and cathode directly facing each other and without intentionally added electrolyte. The maximum chemical yield of 98% of pure product without workup could be obtained with a combination of a glassy carbon anode and a platinum cathode in a single pass of furan solution [178]. Anodic oxidation of benzofuran at a platinum foil in acetonitrile containing lithium perchlorate in the presence of a soluble base such as 2,6-lutidine led, in any case, to a rearranged lactone 3,3-diaryl-2(3H)-benzofuranone [179]. Ar

Ar Ar

–3e–/molecule Y

© 2016 by Taylor & Francis Group, LLC

Ar O

Y

O

O 17–31%

1343

Heterocyclic Compounds

Electrolysis of benzofuran in a CH2Cl2/H2O (1/1) two phase solution in the presence of NaBr afforded 2,3-dibromo-2,3-dihydrobenzofuran (77–84%). Electrolysis of benzofuran in AcOH/H2O (100/1) in the presence of NH4Br, regioselective bromination at the C(5)-position of benzofuran took place to afford 5-bromo-benzofuran (48%) [180]. Electrocatalytic oxidation of 2,5-bis-dihydroxymethylfuran in alkaline medium at a platinum electrode modified by lead adatoms leads to 80% furan-2,5-dicarbaldehyde [181]. Anodic fluorination of 3-methylbenzofuran and ethyl(3-benzofuranyl)acetate was successfully carried out to give mainly 2,3-difluoro-2,3-dihydrobenzofuran derivatives employing Et4NF · 4HF and Et3N·3HF [182]. 2. Thiophene and benzo Derivatives Electroreduction of polyhalothiophenes results in the preferential reduction of the α-halogen atoms. Preparative electrochemical reduction of these products at controlled potential can be used to synthesize 2,3,4-trihalothiophenes, 3,4-dihalothiophenes, and 3-halothiophenes. Depending on the experimental conditions, good yield on the desired product can be obtained [183]. When 2,5-dibromothiophene is reduced, at potentials corresponding to its first voltammetry wave, the main product is the expected 2-bromothiophene. An induced halogen dance can take place, as side process, producing 2,4- and 2,3-dibromothiophenes [184]. The reaction of thiophene derivatives with an electrochemically generated diarylcarbenium ion pool afforded multiple alkylation products [184a]. Preparative cathodic reduction of benzo[b]thiophene at a constant current in aqueous and mixed organic–aqueous solution of TBA salts, at a mercury pool as cathode, led to 2-ethylbenzenethiol [185].

–

C2H5 +H+

+2e– +H+ S

SH

S–

S

+2e– +2H+

SH

The electrooxidation of 2,5-dimethyl- and tetramethyl-thiophene in MeOH containing sodium methoxide produced isomeric mixtures of the corresponding 2,5-dimethoxy adducts, together with side-chain oxidation products. When sodium acetate was used as electrolyte, the products were the corresponding 2-(methoxymethyl)thiophenes and thiophene-2-carbaldehydes [186]. Anodic oxidation of readily available 4- or 7-methoxybenzo[b]thiophenes affords the bisketals of benzo[b]thiophene-4,7-quinones in excellent yields [187]. Anodic fluorination of 3-oxo-2,3-dihydrobenzo[b]thiophene and methyl 3-oxo-2,3dihydrobenzo[b]thiophene-2-carboxylate gave the corresponding monofluorinated products selectively in moderate yields [187a].

D.

COMPOUNDS CONTAINING ONE NITROGEN AND ONE OXYGEN/SULFUR ATOM IN THE RING

1. Isoxazoles and benzoxazoles A quaternized isoxazole can be ring-opened by reduction to provide a derivative of a β-diketone in good yield [188]. R O

NHMe

N

´R

+

O

© 2016 by Taylor & Francis Group, LLC

Me

pH = 5–6

´R

O

H2O, H+

+2e–, +H+ R

´R

O

R

1344

Organic Electrochemistry

Reductive electrolysis of 5-substituted isoxazoles in nonaqueous media on a cathode with high (graphite) and low (ordinary steel) hydrogen overvoltage is accompanied by proton removal from position 3 of the isoxazole ring, which are readily isomerized to linear enolate anions. Acidification of the final reaction mixture during the standard workup results in β-cyanoketones as the major electrolysis products [189]. Preparative electrolysis of 2-(bromodifluoromethyl)benzoxazole in anhydrous DMF–TEABF4 at −1.45 V (vs. SCE) on a carbon felt cathode gave 2-(difluoromethyl)benzoxazole in 65–70% yield [190]. The electrochemical deprotonation of benzoxazolone in an ionic liquid (1-butyl-3-methyl imidazolium, BMIM–BF4), followed by the addition of ethyl iodide, led to the isolation of the corresponding 5-chloro-ethylbenzo[d]oxazol-2(3H)-one in a very high yield [191]. 2. Thiazoles Electroreduction of thiazole derivatives proceeded to provide the corresponding 2-mercaptothiazoles in good yield. Although the substrates contain other reducible groups such as carbonyl, nitrile, and ester groups, only reductive cleavage of the C–S bond took place selectively [192]. H2N

H2N

N

N

+2e–, +2H+ Y S

S

Y

–MeY

SH

Y S

Y = PhCo 69% Y = CN 53% Y = EtOCO 70%

Anodic fluorination of 2-(4-phenylthiazolyl)propargyl sulfide was carried out mainly at constant potential (+1.6 V vs. SCE) in Et3N·5HF/dimethoxyethane to give the corresponding monofluorinated thiazole and trifluorinated thiazoline derivative [193]. N

N

N S

S

E.

Ph

Ph

Ph Divided cell –e–

F

+ F F

S

S

F S

S

COMPOUNDS CONTAINING TWO NITROGEN AND ONE SULFUR ATOM IN THE RING

1. Thiadiazole and Derivatives The electrochemical reduction of 3-halogen-1,2,4-thiadiazole-5-sulfenylchlorides led to substituted cyanimidodithiocarbonates [194]. SR

SR

Cl N +2e– – SR –Cl

N

+Mel N

N

S

N

N

S Me

–

S

VI. A.

ELECTROSyNTHESIS OF FUSED HETEROCyCLIC COMPOUNDS COMPOUNDS CONTAINING TWO NITROGEN ATOMS IN THE MOLECULE

1. Naphthyridines and Derivatives The cathodic reduction of 3,5-dicyano-1,4-dihydro-2,6-dimethyl-4-(o-nitrophenyl) pyridine in protic medium, at a mercury cathode, leads quantitatively to 5-amino-3,10b-dihydrobenzo[c][2,7] naphthyridine 6-oxide. The N-oxide group can be either reduced at –1.4 V versus SCE in the same medium or anodically oxidized at a graphite felt anode [195].

© 2016 by Taylor & Francis Group, LLC

1345

Heterocyclic Compounds

NC

O

NHOH

NO2 NC

CN

N NC

CN

NH2

–0.6 V (vs. SCE) +4e–/+4H+ Me

Me

Me

N H

O

N H

NH2

Me

Me

+2e–/+2H+

Me

N

Me

N H

–H2O

NH2 N

Me

+2e–/–2H+ N

NH

Me

NC

Me

NC

When the cyano groups in 3 and 5 positions of the substrate are changed to methoxycarbonyl groups, 6-hydroxy-1-methoxycarbonyl-2,4-dimethyl-5-oxobenzo[c][2,7]naphthyridine (nifedipine) is obtained [196], and when only the cyano group in 5 position is replaced by a methoxycarbonyl group, 1-cyano-6-hydroxy-2,4-dimethyl-5-oxobenzo[c][2,7]naphthyridine is the product [197]. Naphthyridine rings have been prepared by electrochemically induced hetero[4+2] cycloaddition reactions between 2-vinylpyrroles and β-acceptor enamines [198,199]. 2. Quinazolines Indolo[2,1-b]quinazoline-6,12-dione (tryptanthrin) has been obtained [200] in high yield by cathodic reduction of isatin in dichloromethane/Et4NCl at –0.84 V (vs. Ag/Ag+) employing a Hg, Pt, or Pb cathode and with a charge consumption corresponding to a 0.5 electron/substrate-molecule process. In this reaction, an electron transfer to the oxygen in air is involved. O

O

O

O O NH

O

1

+e– –H

+1 O

N H

N

N –

–O

a

O

O–

O O NH

a

+

O2

O

N

Workup

+ O 2–

O

O N

COOH N O

O N

–CO2 O

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–H2O

O

O N

O NH O O

O–

HN

O H2O

O

O N

O O

NH

1346

Organic Electrochemistry

3. Pyrrolopyrazine In slightly basic aqueous EtOH medium or in acetonitrile, electrochemical reduction of 4-methyl-5(2-pyrazinyl)-1,2-dithiole-3-thione affords a convenient route to pyrrolo[1,2-a]pyrazone derivatives, which are found as metabolites of the drug in host urine [201]. N

S

N S

S

+2e–

N

N

S

Me

−S

Me

−

S

−

−

–S2– Fast

+2e–

N

N

Me

Me

N +

N S−

S S

S

SH +H+

HN Me N S

4. Pyridoimidazole The electrochemical reduction of N-(2-nitro-4-R-phenyl)pyridinium chlorides at a lead cathode in HCl/iPrOH/H2O affords the corresponding 7-R-pyrido[1,2-a]benzimidazoles in 70–94% yield [201a]. + N

R

Cl –

4F

N

Pb cathode

R

iPrOH, 4% HCI

O2N

B.

N

COMPOUNDS CONTAINING THREE NITROGEN ATOMS IN THE MOLECULE

1. Pyrimidoindol Derivatives The electrochemical oxidation of catechol in the presence of 6-amino-2,3-dihydro-2-thioxopyrimidin-4(1H)-one as nucleophile at carbon rod electrodes in an undivided cell under controlled potential conditions affords 2,3-dihydro-6,7-dihydroxy-2-thioxo-1H-pyrimido[4,5-b]indol-4(9H)-one, an uracil derivative [202]. NH2

H N

HN

OH

S

OH

HN

HN O

OH

S HN

OH

O

C.

COMPOUNDS CONTAINING THREE NITROGEN AND ONE OXYGEN ATOMS IN THE MOLECULE

1. Imidazo-Oxadiazines Cathodic reduction of phenacyl bromide semicarbazones in DMF/LiClO4 under high diluted conditions on a mercury cathode afforded in very good yield 3,7-diaryl-2H-imidazo[2,1-b]-1,3,4oxadiazines [203]. N 2 Ar

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NH-CONH2 Br

N +2e– –0.8 V (vs. SCE)

N Ar N

O

Ar

1347

Heterocyclic Compounds

The formation of this imidazo-oxadiazine can be explained as follows:

N

H N CO NH2

N Ar

CH2Br

H N CO NH2

–N

–Br–

+ 2e–

Ar

H N CO NH2 Ar

CH2–

CH2

1

1 CH2 H

H

CH2

Ar

+ N N

NH2 N

N

Ar

Ar

Ar

N

N

Ar

O

O

H N

–Br–

CO

N

–NH3

N H

N O–

H N CO NH2 NH2

Ar Br

–H+ N

Ar

N Ar N

O

Cathodic reduction of N-2-substituted benzimidazolyl imidate in DMF/Bu4NClO4 afforded 1,3,5-oxadiazino[1,2-a] benzimidazol [204]. N

N OEt

Cathodic reduction

N

N N

N OEt

O

D.

72%

O

EtO

COMPOUNDS CONTAINING THREE NITROGEN AND ONE SULFUR ATOM IN THE MOLECULE

1. Thiazolotriazines The electrochemical oxidation of catechols in the presence of 6-methyl-1,2,4-triazine-3-thion-5-one in aqueous sodium acetate produces 7H-thiazolo[3,2-b]-1,2,4-triazin-7-one derivatives through a Michael-type addition reaction to the anodically generated o-quinones [205]. OH

OH

E.

O

N

Me

N

N

O

–2e–, –2H+ S–

S NH

Me

N

NH

OH OH

O –2e– –2H+

N

S

OH

N

Me

N

OH 74%

COMPOUNDS CONTAINING FOUR NITROGEN AND ONE SULFUR ATOM IN THE MOLECULE

1. Triazinothiadiazines Electrochemical oxidation of catechols in the presence of 4-amino-6-methyl-1,2,4-triazine-3-thion5-one as a nucleophile in aqueous solution led to the efficient synthesis of 1,2,4-triazino[3,4-b]1,3,4-thiadiazines [206]. O

NH2

R OH

O

N

SH –2e–, –2H+

+ OH

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Me

N N H

H N

Me

R OH

N N N

S R = H 71% R = OH 65%

OH

1348

Organic Electrochemistry

Anodic formation of several 3,6-disubstituted 1,2,4-triazolo[3,4-b][1,3,4]thiadiazoles was achieved from aryl aldehyde N-(5-aryl-1,3,4-thiadiazol-2-yl) hydrazones [126].

VII.

ELECTROCHEMISTRy OF FUSED HETEROCyCLIC COMPOUNDS

A. COMPOUNDS CONTAINING THREE NITROGEN ATOMS IN THE MOLECULE 1. Pyrazolopyrimidines The pyrazolo[1,5-a]pyrimidine-7-amines are electrochemically reduced in hydroorganic medium at low pH into the corresponding 4,5-dihydrocompounds through a deaminative reduction step [207]. NH2 R6

R6

N

N

R2 R5

N

+2e–, +2H+

R2

–NH3

R5

N

N N

+2e–, +2H+

R2 R5

N R3

R3

B.

R6

N

N H

R3

COMPOUNDS CONTAINING FOUR NITROGEN ATOMS IN THE MOLECULE

1. Purines and Derivatives Purines (adenine and other 6-substituted derivatives) are irreversibly reduced in aqueous media, on mercury electrodes, either in two successive steps, where each step involves the addition of two electrons and two protons to form the hydrogenated species or in a single 4e– reduction involving the hydrogenation of the two N═C bonds. In the case of adenine, relatively slow elimination of ammonia from the reduced form has also been described [208]. Direct guanine reduction, in ionic liquids as solvents, takes place via one-electron process to give an anion radical, which is further protonated by the cations of the solvent. The obtained radical undergoes final dimerization [209]. O H

H2N

O N

N N

N H

+e–, +H+ Ionic liquid

H

H2N

N N

H N N H

O H Dimerization H2N

N N

H N

H N

N H

N H

N

NH2 N

H

O

Electrochemists have done much work on the electrochemical reactions of nucleic acids, indicating that bases in the nucleic acid are reduced and/or oxidized at electrode surface. Attempts to highlight how the DNA/electrode interface can be used to design electrochemical assays for nucleic acids detection (biosensors) have been reported [210]. The electrooxidation of 3,7-dimethylxanthine at solid electrodes proceeds in a single 4e −/4H+ pH-dependent process to give a diimine species, which decomposes in chemical follow-up steps [211]. The electrochemical oxidation of various N-methylated uric acids has been studied at a pyrolytic graphite electrode at physiological pH ═ 7.2. The Ep value was found to shift to less positive potentials when a methyl group is present at the N−1 position and to more positive potentials when substitution is at the N−3 position or at nitrogens of the imidazole ring [212].

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1349

Heterocyclic Compounds O

O

O H3C

H N

N

–2e–, –2H+

H3C

N

N

O O

N H

H2O

O +2e–, +2H+

N

O

N

OH H N

N

O

OH–

N

CH3

H3C O

N OH

N H

CH3

CH3

pH = 7 O

OH H N

HN

NH O

Me

O

+ O

N

O

H N

O

Me

N

N H

CH3

H

CH3

Cyclic voltammetry of various purine derivatives, such as aminophylline, caffeine, and theophylline, shows that they had a similar behavior under the same conditions. The best electrodic material in the oxidation field of potential is electrochemically activated glassy carbon [213]. The electrochemical oxidation of theophylline [214] and caffeine [215] in aqueous medium might be expressed as O H 3C

N

O

N

O H N N

–2e–, –2H+ +2H2O

H3C

N

O

N

O

OH H N

H3C

N OH

O

CH3

CH3

N

N N

O

CH3

N

–2e–, –2H+ +H2O –HCHO

H 3C

H N

N

O

N

N CH3

CH3

Electrochemical oxidation of xanthosine at pH ═ 2 at a pyrolytic graphite electrode generates an electrophilic radical cation intermediate. Nucleophilic attack by xanthosine on this radical results ultimately in the formation of 3-(8-xanthosyl)xanthosine. Hydrolytic cleavage of one ribose residue in acidic solution leads to 3-(8-xanthosyl)xanthine [216]. O H

O

N

N H R = Ribose

H–

O N

H –e–

N R

O

N

N N H

+ N R

O

O N

N N H

–e–, –2H+

N R

O

N

N

N

H

O

O H

O

N N H

N R

R N

N

The electrochemical oxidation of guanosine passes through an 8-oxo derivative intermediate under any condition, thus mimicking the primary effects of oxidative stress in biological systems. The 8-oxoderivative is immediately oxidized at potentials where the oxidation of guanosine occurs to an intermediate species with di-imine structure [217,218]. Oxidation of guanosine with electrochemically generated superoxide (O2−•) leads to the imidazole derivative as a single-electron oxidation product of guanosine. A crucial step in the mechanism of the oxidation is the proton-coupled electron transfer from guanosine to the hydroperoxy radical (HO2•) that is derived from O2−• [219]. The oxidation of N9-substituted adenine on pyrolytic graphite electrodes takes place at the adenine moiety, yielding adsorbed compounds with a common structure that act as a very efficient catalyst of the oxidation of NADH at low potentials [220].

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1350

Organic Electrochemistry

2. Pteridines Electrochemically pteridines behave as quinoxalines, the pyrazine ring being reduced [221]. Systematic electrochemical studies with leucopterin over a wide pH range indicate that there is an interesting pH dependence of the reduction process, leading to an isoxanthopterin derivative at low pH and a xanthopterin derivative at higher pH [222]. O

O N NH

NH

+4e

O NH

O

NH

H+ H 2N

NH

N

H2 N

O

NH

+4e N

NH

O

+OH–

O

H2N

N

N

C. COMPOUNDS CONTAINING TWO NITROGEN AND ONE OXYGEN ATOM IN THE MOLECULE 1. Pyridooxazines Regioselective anodic fluorination of 2H-1,4-pyrido[3,2-b]-1,4-oxazin-3(4H)-one derivatives was successfully carried out in dimethoxyethane containing Et4NF · 4HF using an undivided cell to provide the corresponding α-monofluorinated products [223]. R N

N

N

O –2e–, –H+

R N

O

O

F

Et4NF∙4HF O

Selective electrochemical fluorodesulfurization of benzo- and pyrido-fused oxazine derivatives using ex-cell halogen mediators has been achieved [223a].

D.

COMPOUNDS CONTAINING TWO NITROGEN AND ONE SULFUR ATOM IN THE MOLECULE

1. Thienopyrazines Cathodic reductions of thieno[2,3-b]pyrazines and thieno[3,4-b]pyrazines both lead, in aqueous medium, to a dehydro compound where the two nitrogen atoms of the pyrazine ring are hydrogenated [224].

E.

COMPOUNDS CONTAINING THREE NITROGEN AND ONE OXYGEN ATOM IN THE MOLECULE

1. Imidazo-Oxadiazines 4-Aryl-1-(1-arylethylideneamino)-1,3-dihydro-2-imidazolones and 4-aryl-1-(1-aryl-ethyl amino)1,3-dihydro-2-imidazolones have been selectively obtained by controlled potential reduction (at −1.6 or −1.9 V, respectively) of 3,7-diaryl-2H-imidazo[2,1-b]-1,3,4-oxadiazines [225]. N

N

Ar

Ar

N

+2e +2H+

N

Ar

Ar N

EtOH/LiClO4

N

O

O

H NH

Ar

N

Ar

Ar N H

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NH

O

+

EtOH/LiClO4 +2e +2H+

N H

O

1351

Heterocyclic Compounds

F. COMPOUNDS CONTAINING FOUR NITROGEN AND ONE SULFUR ATOM IN THE MOLECULE 1. Triazolothiadiazines Constant potential anodic oxidation of s-triazolo[3,4-b]-1,3,4-thiadiazine derivatives in dimethoxyethane containing Et4NF · 4HF, using an undivided cell, provided the corresponding 7-monofluorinated products. 7,7-Difluorination was also anodically achieved [226]. S

N

S

N –2e–/–1H+

F –2e–/–1H+

N

N N

N

Ph

R

N

R

N N

Ph

F F

N

N N

S

R Ph

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Heterocyclic Compounds

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113. 114. 115. 116. 117. 118. 119. 120. 121. 122. 123. 124. 125.

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126. 127. 128. 129. 130. 131. 131a. 131b. 132. 133. 134. 135. 136. 137. 138. 139. 140. 140a. 141. 142. 143. 144. 145. 146. 147. 148. 149. 149a. 150. 151. 152. 153. 154. 155.

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Heterocyclic Compounds 156. 157. 158. 159. 160. 161. 162. 163. 164. 165. 166. 166a. 166b. 166c. 167. 168. 169. 170. 171. 172. 172a. 172b. 173. 174. 175. 176. 177. 178. 179. 180. 181. 182. 183. 184. 184a. 185. 186. 187. 187a. 188. 189. 190. 191. 192. 193. 194. 195. 196. 197. 198. 199. 200.

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35

Organoelemental Compounds Jun-ichi Yoshida, Toshiki Nokami, and Seiji Suga

CONTENTS I. II.

Introduction ........................................................................................................................ 1357 Electrochemical Synthesis of Organoelemental Compounds............................................. 1358 A. Oxidative Formation of Element–Carbon Bonds ....................................................... 1358 1. Oxidation of Organoelemental Compounds Using Reactive Anodes.................. 1358 B. Reductive Formation of Element–Carbon Bonds ....................................................... 1359 1. Reduction of Organic Compounds Using Reactive Cathodes ............................. 1359 2. Reduction of Organic Compounds in the Presence of Element Halides ............. 1360 3. Reduction of Organic Compounds Using Reactive Anodes of an Element......... 1363 III. Electrochemical Oxidation of Organoelemental Compounds ............................................ 1363 A. Electrochemical Oxidation of Group 1 and 2 Organoelemental Compounds ............ 1363 B. Electrochemical Oxidation of Group 13 Organoelemental Compounds .................... 1364 C. Electrochemical Oxidation of Group 14 Organoelemental Compounds .................... 1365 1. Electrochemical Oxidation of Tetraorganoelemental Compounds ...................... 1365 2. Electrochemical Oxidation of Dielemental Compounds (R3E–ER3) .................. 1376 3. Electrochemical Oxidation of Element–Element Double Bonds (R2E═ER2, R2E═ER–ER═ER2).......................................................................... 1377 4. Electrochemical Oxidation of Organoelemental Hydrides (R3EH)..................... 1379 5. Electrochemical Oxidation of Stable Radicals of the Heavy Group 14 . Elements (R3E ) ................................................................................................... 1379 6. Electrochemical Oxidation of Silylenes and Germylenes ((R2N)2E:) .................. 1380 7. Electrochemical Oxidation of Other Types of Group 14 Organoelemental Compounds .......................................................................................................... 1380 IV. Electrochemical Reduction of Organoelemental Compounds ........................................... 1381 A. Electrochemical Reduction of Group 13 Organoelemental Compounds.................... 1382 B. Electrochemical Reduction of Group 14 Organoelement Halides .............................. 1382 1. Reduction Potentials and General Reaction Patterns of Group 14 Organoelement Halides ....................................................................................... 1382 2. Reductive Formation of Element–Element Bonds ............................................... 1383 3. Reductive Formation of Element–Carbon Bonds ................................................ 1385 C. Electrochemical Reduction of Group 14 Organoelemental Compounds Containing Reducible π- and σ-Systems........................................................................................ 1386 D. Electrochemical Reduction of Group 15 Organoelemental Compounds.................... 1386 References .................................................................................................................................... 1387

I. INTRODUCTION The organoelemental compounds such as group 1, 2, 13, and 14 organoelemental compounds have played a major role in organic synthesis and enjoy widespread use in both laboratories and industry. Some organoelemental compounds such as organosilicon compounds are also utilized as materials. Some organoelemental compounds, such as group 1 and 2 organoelemental compounds, are also called organometallic compounds. However, in this chapter, we call them organoelemental compounds. 1357

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The electrochemistry of other organometallic compounds, especially transition metal compounds, is discussed in Chapter 36. This chapter will focus on the electrochemical synthesis and reactions of organic compounds containing main group elements. The electrochemistry of organoelemental compounds has a long history, and enormous advances have been made in the study of this field so far. The scope of this subject is, therefore, too extensive to be covered completely in a single chapter. There are excellent review articles on general electrochemistry of organoelemental compounds [1–9] and specific topics [10–12] in this field. Thus, the approach here is to present those fundamental aspects of electrochemical synthesis and reactions of organoelemental compounds with special emphasis on recent developments. This chapter is not an exhaustive compilation of all known reactions, but rather sampling of sufficient variety to illustrate the fascinating chemistry of electrochemical synthesis and reactions of organoelemental compounds.

II.

ELECTROCHEMICAL SyNTHESIS OF ORgANOELEMENTAL COMPOUNDS

Electrochemical reactions serve as efficient and convenient methods for the synthesis of organoelemental compounds. There are four major methods for the formation of element–carbon bonds. The first method utilizes the anodic oxidation of organoelemental compounds using reactive anodes. In the second method, the organic compounds are reduced using reactive cathodes of the elements. The third method involves the cathodic reduction of organic compounds in the presence of element halides. The fourth one utilizes both the cathodic and anodic processes.

A. OXIDATIVE FORMATION OF ELEMENT–CARBON BONDS 1. Oxidation of Organoelemental Compounds Using Reactive Anodes Organoelemental compounds can be synthesized by anodic reactions. This reaction proceeds by the initial oxidation of the starting organoelemental compounds (R–E1) to generate an organic radical (R•), which reacts with a reactive anode of an element (E2) to form the desired product (R–E2) (Equation 35.1). In some cases, a mechanism involving the ionic reaction of the cation of E2 formed at the anode surface with R–E1 is suggested. R–E1

–e–

+

R

+

E1

E2 (anode)

(35.1)

R–E2

The anodic oxidation of Grignard reagents using reactive anodes of elements such as Al [13,14], Zn [13,15], Mn [15], Cd [13,15], Bi [15], and B [16] is known to give the corresponding organoelemental compounds (Equation 35.2). This reaction has been used for the synthesis of tetraalkyllead in industrial scale (NALCO process) [17,18]. RMgX

+

E (anode)

–e–

RnE

(35.2)

E = Al, Zn, Mn, Cd, Bi, B, Pb (NALCO process)

The anodic oxidation of zinc-, aluminum-, and boron-ate complexes such as Na[ZnR3], Na[AlR4], and Na[BR4] also leads to the formation of the element–carbon bonds (Equation 35.3). Mg [19], Zn [20,21], Cd [21], Hg [21–23], Sn [21,22,24], Pb [21–23,25,26], Sb [21–23], and Bi [21,23] are effective as reactive anodes.

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1359

Organoelemental Compounds

Na[E1Rn]

–e–

E2 (anode)

+

(35.3)

RnE2

E1 = Zn, Al, B E2 = Mg, Zn, Cd, Hg, Sn, Pb, Sb, Bi

B. REDUCTIVE FORMATION OF ELEMENT–CARBON BONDS 1. Reduction of Organic Compounds Using Reactive Cathodes Reduction of some organic compounds at reactive cathodes leads to the formation of organoelemental compounds as shown in Equation 35.4. This reaction seems to proceed by one-electron reduction to form radical intermediates, which react with elements to form the element–carbon bonds. +e–

Organic substrate

E (cathode)

(35.4)

R–E

R

Organic halides serve as effective precursors of organic radicals, and the reduction of organic halides with reactive cathodes of elements leads to the formation of organoelemental compounds (Equation 35.5). For example, the reduction of alkyl halides with a Pb cathode results in the formation of R4Pb [27–31]. The reduction with a Hg cathode results in the formation of R2Hg [32–35]. Sn is also effective as a reactive cathode to give R4Sn [36–38]. +

R–X

+e–

E (cathode) E = Pb, Hg, Sn

(35.5)

R–E

Carbonyl compounds are also used as starting materials for the reductive synthesis of organoelemental compounds. The electrochemical reduction of carbonyl compounds such as ketones and aldehydes using reactive cathodes composed of the respective elements gives the deoxygenated organoelemental compounds although yields and selectivities are generally not high (Equation 35.6). Ge [39], Pb [40,41], Sn [42], Hg [41,43], and Zn [44] are known to be effective as reactive cathodes. The mechanism involving α-hydroxy carbon radicals is suggested [45]. R O

+

E (cathode)

R΄

+e– –H2O

R

R

(35.6)

E R΄

R΄

Activated olefins such as acrylonitrile and methyl vinyl ketone are also effective as precursors of organoelemental compounds. The electrochemical reduction of activated olefins using Sn [46,47], Bi [48], Se [49], Te [49], and Hg [50] cathodes gives the corresponding organoelemental compounds as shown in the following equation. Y

+

E (cathode)

+e– +H+

Y n

(35.7)

E

Y = CN, COMe, COOH E = Sn, Bi, Se, Te, Hg

Diaryliodonium ions are also reducible and their cathodic reduction using a Hg cathode affords Ar2Hg (Equation 35.8) [51]. Benzylsulfonium ion is also effective as a substrate [52]. I+OH–

X 2

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+

Hg (cathode)

+e– NaOH–H2O

X

Hg 2

(35.8)

1360

Organic Electrochemistry

2. Reduction of Organic Compounds in the Presence of Element Halides The combination of organic halides and element halides is also effective for the reductive formation of the element–carbon bond. This type of reaction proceeds by a mechanism involving the cathodic reduction of organic halides (Equation 35.9) because reduction potentials of organic halides are usually less negative than those of element halides. The carbanion intermediate is usually responsible for the formation of organoelemental compounds. +2e– R

X

–X–

R–

E–X R

(35.9)

E

Several examples have been reported for the synthesis of organosilicon compounds. The reduction of organic halides such as allyl, benzyl, aryl, and alkenyl halides in the presence of chlorosilanes in a divided cell gives the corresponding organosilicon compounds (Equation 35.10) [53–55]. It is interesting that the regioselectivity of the reaction of allyl halides depends on the nature of the silylating agent. Trimethylsilyl and dimethylphenylsilyl groups are introduced to the less substituted end of the allyl group, whereas the dimethylsilyl group is introduced to both ends of the allyl group.

Ph

Cl +

+2e– (2.43–2.85 F)

R3SiCl

Ph

SiR3

Et4NOTs/DMF

+

Ph

(35.10) SiR3

R3Si

Ratio

Yield

Me3Si

100:0

98

PhMe2Si

100:0

66

HMe2Si

50:50

84

Palladium-catalyzed reductive silylation has also been reported. The cathodic reduction of allylic acetates in the presence of silylating agents and a catalytic amount of (Ph3P)4Pd [56] gives the corresponding allylsilanes. The initially formed π-allylpalladium(II) complex seems to be reduced at the cathode to generate the allyl anion intermediate, which reacts with chlorosilane to give the final product. The Al-promoted alkylation of SiCl4 with organic halides followed by the electrochemical regeneration of Al is also noteworthy [57]. This reaction serves as an effective method for the synthesis of tetraalkylsilanes from the corresponding alkyl halides in a large scale. Reductive silylation of organic halides is effectively accomplished using a sacrificial Al or Mg anode and hexamethylphosphoramide (HMPA) as a cosolvent as shown in Equation 35.11 [58–61]. It is advantageous that this reaction can be carried out in an undivided cell. It is also worth noting that aryl chlorides, which are less easily reduced than the corresponding bromides and iodides, can be used as organic halides. Cl +

Me3SiCl

+2e– (2.2 F) Al anode

SiMe3

(35.11)

THF/HMPA (80:20) 75%

Reductive silylation of mono- and polyhalothiophenes has also been achieved using an Al sacrificial anode (Equation 35.12) [62–64]. This reaction provides a convenient method for the preparation of 2,5-bis(trimethylsilyl)thiophene that serves as a good precursor of polythiophene. The sacrificial Mg and Al anode technique was also successfully applied to the reductive silylation of bromopyrroles [65].

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1361

Organoelemental Compounds +4e– (4.4 F) Al anode Cl

Cl

S

+ Me3SiCl Bu4NBr THF–HMPA (3:1)

Me3Si

SiMe3

S

(35.12)

93%

Selective mono- and disilylation of polychloromethanes such as carbon tetrachloride and chloroform have been achieved using Zn and Mg sacrificial anodes (Equations 35.13 and 35.14) [66]. The product selectivity seems to depend on the electrode material. The application of this method for the reduction of ClCH2SiMe2Cl leads to the formation of polycarbosilanes [67].

CCl4 + Me3SiCl

+2e– (2.2 F ) Zn anode Et4NBF4 /DMF

(35.13)

Me3SiCCl3 + (Me3Si)2CCl2 94%

66%

+4e– (4.4 F) Mg anode CCl4 + Me3SiCl

Et4NBF4 /DMF

(35.14)

(Me3Si)2CCl2 68%

Fluorine-substituted alkenyl bromides and chlorides were also found to be effective as organic halides [68]. Fluorovinylsilanes were obtained by the electrochemical reduction of the corresponding fluoro-substituted alkenyl chlorides and bromides in the presence of Me3SiCl using a sacrificial Al or Zn anode. Activated olefins are also effective for the reductive element–carbon bond formation. The electrochemical reduction of α,β-unsaturated esters, nitriles, and ketones in the presence of Me3SiCl using a sacrificial Mg, Zn, or Al anode affords the corresponding silylated compounds as shown in the following equation [69]. +2e– (2.0 F ) Mg anode

MeO +

Me3SiCl

CO2Et

MeO

(35.15)

Bu4NBr DMF

CO2Et SiMe3 79%

Acyl-elemental compounds are also synthesized electrochemically. The cathodic reduction of acylimidazoles in the presence of Me3SiCl using a Pt cathode and a Pt anode gives the corresponding acylsilanes (Equation 35.16) [70]. A mechanism involving the initial N-silylation of the imidazole moiety with Me3SiCl, which decreases the reduction potential of the amide group, is suggested. The reductive silylation of acylsilane thus produced also takes place under appropriate conditions to give the silyl ether of the α,α-disilyl alcohol. Acid anhydrides are also effective for the reductive silylation. The cathodic reduction of aromatic acid anhydrides in the presence of Me3SiCl also gives the acylsilanes together with the simply reduced alcohols [71]. O

O C7H15

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+

N N

Me3SiCl

+2e– (2.5 F) Bu4NClO4 THF

C7H15

SiMe3 73%

(35.16)

1362

Organic Electrochemistry

Organotin compounds are also synthesized by the cathodic reduction of organic compounds in the presence of tin halides. For example, the reduction of allylic halides in the presence of chlorostannanes gives the corresponding allylstannanes in good yields [72]. Combination of this reaction with in situ palladium-catalyzed reaction with allylic halides leads to effective formation of the headto-tail homocoupling products as shown in the following equation. Cl +

Ph

+2e– Et4NOTs DMF

Bu3SnCl

Ph

SnBu3

Ph

Cl

Ph

(35.17)

2 Ph3P/PdCl2

Ph 89% (1.14 F)

Organoboron compounds are also synthesized by the cathodic reduction of organic compounds in the presence of boron compounds. For example, electrochemical reduction of benzylic halides in the presence of trialkylborates followed by hydrolysis gives the corresponding benzyl boronic acids [73]. The use of pinacolborane as a borating agent leads to the formation of benzylboronic pinacol esters (Equation 35.18). The reactions are carried out in THF using an Mg anode in a single-compartment cell. The electrochemical reduction of dihalogenated compounds in the presence of pinacolborane gives monohalogenated arylboronic esters, although the increase in the amount of electricity leads to the formation of the diboronic esters (Equation 35.19) [74]. The use of allyl halides affords allylboronic pinacol esters as shown in Equation 35.20 [75]. In this case, an Al anode is used. MeO

Cl

+2e– Mg anode

O H

+

B

MeO

B

THF

O

O

O

(35.18)

80% Br Br +2e– Mg anode

O + Br

H

B O

O

B

THF

(35.19)

O Conversion yield 67% +2e– Al anode

O Br +

H

B

THF

O

B (E/Z = 70/30)

O

(35.20)

+

O

B O 81% (90:10)

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O

1363

Organoelemental Compounds

3. Reduction of Organic Compounds Using Reactive Anodes of an Element The combination of the cathodic reduction of organic halides and the reactive anode composed of an element is also effective for the synthesis of organoelemental compounds (Equation 35.21). Examples using Al [76], In [77], and Sn [78–80] as the anode to give the corresponding RAlX2, RInX2, and R2SnX2, respectively, have been reported. The reactions are suggested to proceed by the reduction of RX to generate X−, which is oxidized at the anode to give EX. Organic sulfates are also used for this purpose. For example, the electrolysis of diethylsulfate using an Sn anode in the presence of CH3CH2I gives (CH3CH2)4Sn [81]. RX + E (anode)

Electrolysis

(35.21)

RmEXn

It is worth noting that this method is also effective to produce dielemental compounds. Organodialuminum compounds such as Cl2AlCH2AlCl2 were synthesized from CH2Cl2 [82].

III.

ELECTROCHEMICAL OXIDATION OF ORgANOELEMENTAL COMPOUNDS

A.

ELECTROCHEMICAL OXIDATION OF GROUP 1 AND 2 ORGANOELEMENTAL COMPOUNDS

Electrochemical oxidation of group 1 and 2 organoelemental compounds generally leads to the formation of organic radicals. When a reactive anode of an element is used, the radicals react with the anode to give the corresponding organoelemental compounds (see Section II.A). When an inert anode is used, the radical undergoes coupling, hydrogen abstraction, addition, and disproportionation depending on the conditions as shown in the following equation. R

H

R

R

R

H

SH

R

E

–e–

E+

+

R

Y

+

(35.22) R(-H)

R Y

The electrochemical oxidation of RLi using a Hg anode results in the formation of homocoupling products R–R [83,84]. The electrochemical oxidation of Grignard reagents has been studied extensively [14,85–91]. The anodic oxidation of RMgX also gives the homocoupling product R–R. This process competes with hydrogen abstraction and disproportionation. Addition of the R• to the carbon–carbon double bond takes place when the reaction is carried out in the presence of olefinic compounds such as styrene and butadiene [92]. The electrochemical oxidation of enolates and related compounds has also been studied. For example, the anodic oxidation of Li enolates of esters at low temperatures gives the corresponding coupling products [93]. The anodic oxidation of Na malonates in the presence of olefinic compounds such as vinyl ethers in methanol gives the addition product (Equation 35.23) [94]. The following mechanism is proposed. The one-electron oxidation gives the radical of malonic ester, which adds to the carbon– carbon double bond. The resulting radical is further oxidized to give the cation, which is trapped by methanol to give the final product.

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1364

Organic Electrochemistry CO2Me Na+

CO2Me

–e–

CO2Me

EtO

– CO2Me

CO2Me

EtO

CO2Me

(35.23) +

–e–

MeOH

EtO

B.

CO2Me

MeO

CO2Me

EtO

CO2Me

CO2Me

ELECTROCHEMICAL OXIDATION OF GROUP 13 ORGANOELEMENTAL COMPOUNDS

This section focuses on a rich variety of synthetic applications of the electrochemical oxidation of organoboron compounds, because the electrochemical oxidation of other group 13 organoelemental compounds such as organoaluminum compounds has been studied less extensively. There is a comprehensive review on the electrochemistry of organoboron compounds [95]. In general, the oxidation of triorganoborane (R3B) needs the activation by a nucleophile to form an ate complex (see also Section II.A.1). This negatively charged electron-rich ate complex then undergoes the electron transfer. R3B is readily oxidized at a Pt electrode in methanol containing KOH as a nucleophile to give the homocoupling products (R–R) [96]. Three R groups on B are involved in this reaction; however, yields are not very high. A mechanism involving the initial formation of the borate, which is subjected to one-electron oxidation to form the R• radical, is proposed (Equation 35.24). This mechanism is consistent with a cross-coupling experiment, which gave a nearly statistical distribution of three possible coupling products. The R• radical formed by the anodic oxidation of R4B− anion was detected by a spin-trapping experiment [97]. R3B

R3B–

HO–

+

OH –e–

R3B

OH

+

R

R

R2B

OH

(35.24)

R

The use of a graphite anode, however, gives rise to the formation of different products. The anodic oxidation of R3B using a graphite anode in methanol containing MeONa and NaClO4 gives ROMe as a major product together with the homocoupling product R–R [98]. In this case, the R• radical intermediate undergoes further oxidation at the anode to give the R+ cation, which is trapped by methanol to give ROMe (Equation 35.25). Similarly, the anodic oxidation in acetic acid in the presence of NaOAc leads to the formation of ROAc. The R+ cation also undergoes an elimination reaction to give a regioisomeric mixture of olefins [99,100]. R3B

+

R3B–

MeO–

OMe

–e– R3B

R

OMe

+

R2B

OMe

–

–e

R R+

R–OMe MeOH

© 2016 by Taylor & Francis Group, LLC

R

(35.25)

1365

Organoelemental Compounds

Anodic oxidation of arylboronic esters in the presence of bromide anions leads to the cleavage of the carbon–boron bond to give the corresponding aryl bromides in moderate yields in addition to other products (Equation 35.26) [101]. Product selectivity strongly depends on the structure of the substrate. A mechanism involving the electrochemically generated bromonium ions has been suggested. O MeO

–2e– MeO

B

Br

Et4NBr CH3CN

O

(35.26)

53%

The anodic oxidation of NaBPh4 is also reported to give the coupling product Ph–Ph, but this reaction does not seem to proceed by radical coupling. A mechanism involving the intramolecular carbon–carbon bond formation is proposed [102–104]. Several synthetically useful carbon–carbon bond forming reactions using anodic oxidation of organoboron compounds have been reported. For example, the anodic oxidation of R3B in the presence of butadiene resulted in the transfer of R to butadiene [105]. A mechanism involving the initial formation of R3BOMe−, which undergoes anodic oxidation to generate the R• radical, has been proposed. The anodic oxidation of R3B in the presence of terminal acetylenes to give the coupling products is also interesting [106]. The reaction of R3B with CH3CN is interesting from both mechanistic and synthetic points of view. The electrolysis of R3B in acetonitrile using Bu4NX (X=Br, I) in an undivided cell gives RCH2CN [107]. This reaction is suggested to proceed by the nucleophilic attack of cathodically generated −CH2CN to the anodically generated RI as shown in Equation 35.27. RI seems to be produced by the reaction of R3B and a electrogenerated iodine radical. The electrolysis of R3B in CH3NO2 also seems to proceed in a similar fashion to give RCH2NO2 [108]. Some other electrochemical reactions of R3B have also been reported in the literature [109,110]. I– R3B

–e– I +

I +e–

CH3CN

(35.27)

R–I RCH2CN –CH

2CN

It is notable that BH −4 also undergoes one-electron oxidation to produce H2 and BH3 (B2H6). The BH3 thus formed can be utilized for in situ hydroboration of olefins [111,112].

C.

ELECTROCHEMICAL OXIDATION OF GROUP 14 ORGANOELEMENTAL COMPOUNDS

Group 14 organoelemental compounds, especially organosilicon and -tin compounds, have received significant research interest from viewpoints of electronic structures, reaction mechanism, organic synthesis, and material science. Since the oxidation potentials of tetraorganoelemental compounds especially tetraorganosilanes are very positive, some activation is needed for the anodic oxidation. Although there are several methods for the activation, the use of orbital interactions to raise the HOMO level is commonly utilized for this purpose [113] and various anodic oxidation reactions of organosilicon and -tin compounds have been developed based on this concept. 1. Electrochemical Oxidation of Tetraorganoelemental Compounds a. Unactivated Tetraorganoelemental Compounds The oxidation potential of group 14 tetraorganoelemental compounds increases in the order of Pb 2.5

Ep: Peak potentials determined by cyclic voltammetry using a Pt anode in LiClO4/CH3CN. Source: Yoshida, J. and Nishiwaki, K. J., Chem. Soc. Dalton Trans., 2589, 1998.

+

R3

R3 R1

SiMe3

R1

–e–

R2

R3 R1

SiMe3 R2

R2

R3 –e–

R3 Nu–

R1

R1

R1 Nu +

+ R2

R3

(35.29)

R2

R2

Nu

NuH = ROH, H2O, RNHCO2Me, RNHTs

The oxidation potentials of benzylsilanes are also less positive than those of the corresponding aromatic hydrocarbons and tetraorganosilanes owing to the σ–π interaction. The preparative electrochemical oxidation of benzylsilanes results in the cleavage of the C–Si bond and the introduction of a nucleophile on the carbon [116]. The following example demonstrates that the benzylic C–Si bond is cleaved selectively without affecting the aromatic C–Si bond (Equation 35.30). –2e– C anode undivided cell

Me3Si SiMe3

Me3Si OMe

(35.30)

Et4NOTs/MeOH quantitative

A silyl group exhibits a similar effect for diarylmethanes (Table 35.3) [120–122]. The oxidation potential of diphenyl(trimethylsilyl)methane is much less positive than that of diphenylmethane. Oxidation potentials of diphenylbis(trimethylsilyl)methane and extended diaryl(trimethylsilyl) methanes are slightly higher than that of diphenyl(trimethylsilyl)methane. Selective benzylic C–Si bond cleavage is observed in the anodic oxidation of these compounds. This silyl group effect has been applied to generate and accumulate dendritic diarylcarbenium ions. d. Activation by σ–n Interaction The C–Si σ orbital interacts with a neighboring nonbonding p orbital of heteroatoms such as oxygen, nitrogen, sulfur, and phosphorous (σ–n interaction). Because of the σ–n interaction, the oxidation

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1368

Organic Electrochemistry

TAbLE 35.3 Oxidation Potentials of Diphenylmethane- and Trimethylsilyl-group-Substituted Analogues Ed (V vs. SCE)

Compound

1.81

0.98 SiMe3

1.15 Me3Si

SiMe3 X

X

1.26 (X = H) 1.25 (X = F) X

X SiMe3

Ed: Decomposition potentials determined by rotating disk electrode voltammetry using a glassy carbon anode in LiClO4/CH3CN. Source: Nokami, T. et. al., Tetrahedron, 67, 4664, 2011; Nokami, T. et. al., J. Am. Chem. Soc., 130, 10864, 2008, and unpublished results.

potentials of α-heteroatom-substituted tetraorganosilanes are less positive than those of the corresponding simple heteroatom compounds and tetraorganosilanes [123,124]. Table 35.4 demonstrates that the oxidation potentials of α-heteroatom-substituted organosilicon compounds exhibit less positive oxidation potentials than β-heteroatom-substituted organosilicon compounds, indicating the importance of σ–n interaction. TAbLE 35.4 Oxidation Potentials of Nitrogen-, Sulfur-, and Phosphorous-Substituted Organosilicon Compounds Compound Me3SiCH2NHPh Me3SiCH2CH2NHPh Me3SiCH2SPh Me3SiCH2CH2SPh

Ep/2 (V vs. SCE)

Compound

Ep/2 (V vs. SCE)

0.44 0.60 1.15 1.26

Me3SiCH2PPh2 Me3SiCH2CH2PPh2

0.63 0.88

Source: Cooper, B.E. and Owen, W.J., J. Organomet. Chem., 29, 33, 1971. Ep/2: Half-wave potentials determined by single-sweep voltammetry using a Pt working electrode in Me4NClO4/CH3CN.

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1369

Organoelemental Compounds

TAbLE 35.5 Oxidation Potentials of α-Alkoxysilanes Compound C7H15

Ed (V vs. Ag/AgCl)

Compound

Ed (V vs. Ag/AgCl)

SiMe2Ph 1.33

O

SiMe2Ph

1.95

O

SiMe2Ph

1.39

OMe C7H15

SiMe2Ph

1.54

Me Si Ph 1.64 O Source: Yoshida, J. et al., J. Am. Chem. Soc., 112, 1962, 1990. Ed: Decomposition potentials determined by rotating disk electrode using a glassy carbon anode in LiClO4/CH3CN.

Since the orbital interaction plays a major role, the oxidation potentials of α-heteroatomsubstituted organosilicon compounds depend on the geometry. As a matter of fact, the oxidation potentials of α-alkoxysilanes in which the rotation around the C–O bond is restricted vary dramatically with the torsion angle of Si–C–O–C (Table 35.5) [124]. It is also interesting that the introduction of an additional silyl group on the same carbon causes a further decrease in the oxidation potential probably because of the increase in the population of the favorable conformers. The oxidation potential also depends on the nature of the substituent on silicon (Table 35.6). For example, the introduction of F on silicon causes an anodic shift of the oxidation wave, although the introduction of H on silicon results in a very little anodic shift [125]. Moreover, the oxidation potentials can be tuned by changing substituents of the phenyl group on the silicon atom [126]. The preparative electrochemical oxidation of α-heteroatom-substituted tetraorganosilanes gives rise to the cleavage of the C–Si bond and the introduction of a nucleophile such as methanol on the carbon (Equation 35.31). Therefore, silyl groups serve as electroauxiliaries [9] for electrochemical oxidation of heteroatom compounds. An electroauxiliary is a functional group that activates the

TAbLE 35.6 Oxidation Potentials of Sila-Functionalized α-Heteroatom-Substituted Organosilicon Compounds Compound

Ed (V vs. Ag/AgCl)

menthyl O

SiMe3

1.65

menthyl O

SiMe2H

1.73

menthyl O

SiMe2F

1.81

Compound

Ed (V vs. Ag/AgCl)

SiMe3

1.17

PhS

SiMe2H

1.17

PhS

SiMe2F

1.27

PhS

Source: Yoshida, J. et al., Inorg. Chim. Acta, 220, 129, 1994. Ed: Decomposition potentials determined by rotating disk electrode voltammetry using a glassy carbon anode in LiClO4/CH3CN.

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1370

Organic Electrochemistry

parent molecule toward electron transfer and controls the reaction pathway. Various reactions of this type have been reported for compounds containing oxygen [127–129], nitrogen [130], sulfur [126,129,131–133], and selenium [133]. +

–e– Y

SiMe3

Y

–e–

Y: RO, RNCO2Me, RS, RSe

NuH

Y

SiMe3

Y

+

Nu

(35.31)

NuH: ROH

Several synthetic applications of the electrochemical oxidation of α-heteroatom-substituted organosilicon compounds have been developed. Since the anodic oxidation of α-alkoxysilanes in alcohols leads to the formation of acetals (Equation 35.32), which are readily hydrolyzed to carbonyl compounds such as aldehydes and ketones, α-alkoxysilanes serve as synthons of carbonyl groups [128,129]. A general iterative method for the synthesis of optically active polyols was developed based on this concept [134]. The anodic oxidation of α-phenylthiosilanes in alcohols results in the formation of the corresponding S,O-acetals, which undergo further electrochemical oxidation to give the corresponding acetals (Equation 35.33). Therefore, α-phenylthiosilanes also serve as synthons of carbonyl compounds [132]. The anodic oxidation of α-nitrogen-substituted organosilanes is also useful in organic synthesis [130]. For example, the anodic oxidation of α-silyl carbamates gives α-methoxy carbamates regioselectively (Equation 35.34). The decrease in the oxidation potential is also advantageous. This chemistry was successfully applied for the regioselective functionalization of β-lactams [135,136].

OMe

R

–2e– C anode undivided cell

SPh

Et4NOTs/MeOH

SiMe3

CO2Me R

N R΄ SiMe3

R

Et4NOTs/MeOH

SiMe3

R

–2e– C anode undivided cell

R

–2e– C anode undivided cell Et4NOTs/MeOH

(35.32)

OMe

SPh OMe

OMe

R

–2e– MeOH

OMe

(35.33)

OMe

CO2Me R

N R΄

(35.34)

OMe

α-Heteroatom-substituted organosilicon compounds serve as good precursors of highly reactive N-acyliminium ion pools (Equation 35.35). Low-temperature electrochemical oxidation of α-silyl carbamates affords a solution of the corresponding N-acyliminium ions as a cation pool [137]. Thus, generated N-acyliminium ions react with aromatic compounds such as 1,3,5-trimethoxybenzene to give Friedel–Crafts type alkylation products [138,139]. N-Acyliminium ion pools undergo an inverse electron demand Diels–Alder reaction with alkenes and alkynes to give [4+2] cycloadducts [140,141]. An N-acyliminium ion pool also serves as an initiator for living cationic polymerization [142,143].

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1371

Organoelemental Compounds OMe OMe N MeO

OMe Bu OMe

MeO –2e– C anode divided cell

CO2Me N

CO2Me

N

Bu4NBF4/CH2CI2 –78°C

SiMe3

(35.35)

Bu R

N + Bu

Bu

CO2Me

O

R

O

CO2Me OR

N

OR

Bu

+

n OR

The anodic oxidation of sulfides and ethers having two silyl groups in methanol leads to the formation of the corresponding methyl esters (Equation 35.36). Since (phenylthio) bis(trimethylsilyl)methane and methoxybis(trimethylsilyl)methane can be easily deprotonated and alkylated, these compounds serve as effective synthons of the anion of the methoxycarbonyl group [128,129,132]. –4e– C anode undivided cell

SiMe3 OMe

R

Et4NOTs/MeOH

SiMe3

(35.36)

OMe

R O

The sequential transformation of two silyl groups on the same α-carbon of nitrogen has been achieved under the cation pool conditions (Equation 35.37) [144]. Trimethylsilyl groups act as electroauxiliaries and enable the introduction of two different nucleophiles on the same α-carbon of the nitrogen atom. The method is also applicable to substrates having two trimethylsilyl groups on two different α-carbon atoms [145]. Me3Si R

SiMe3

N

–2e– C anode divided cell

SiMe3

CO2Me

+ N

R

R΄ Bu4NBF4/CH2CI2 –78°C

–

–2e C anode divided cell Bu4NBF4/CH2CI2 –78°C

Me3Si

R΄

R

CO2Me

+ N CO2Me

N

R΄

CO2Me

(35.37)

R1 R

R1

R1M

1 R2 R

R2M R΄

R

N

R΄

M: metal

CO2Me

The activation by σ–n interaction is also effective for other organoelemental compounds. For example, the oxidation potentials of α-heteroatom-substituted tetraorganogermanes are less positive than those of the corresponding organosilicon compounds [146]. It is interesting that the anodic

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oxidation of ethers having a silyl group and a germyl group on the same α-carbon results in the selective cleavage of the C–Ge bond. In the cation radical intermediate, the C–Ge bond seems to be more susceptible toward the cleavage than the C–Si bond. The C–Sn σ orbital also interacts with a neighboring nonbonding p orbital of heteroatoms, and this interaction decreases the oxidation potential significantly [147,148]. It should be noted that the oxidation potentials of α-heteroatom-substituted tetraorganostannanes are less positive than those of carbon nucleophiles such as allylsilanes and enol silyl ethers. Consequently, such carbon nucleophiles can be used for the electrochemical oxidation of α-heteroatom-substituted tetraorganostannanes to achieve carbon–carbon bond formation. In fact, various intermolecular carbon–carbon bond formation reactions have been developed using α-heteroatom-substituted tetraorganostannanes [149–151]. The anodic oxidation of heteroatom compounds having a silyl group and a stannyl group on the same α-carbon has been examined [152]. The anodic oxidation in the presence of the allylsilane led to the selective cleavage of the C–Sn bond and the introduction of the allyl group on the carbon (Equation 35.38). Probably, the C–Sn σ orbital interacts more effectively with the p orbital of the heteroatom than the C–Si σ orbital in the cation radical intermediate. –e– C anode undivided

Y SiMe3

+

SnBu3

R

BU4NCIO4 CH2CI2

SiMe3

Y R

(35.38) SiMe3

Y = R΄O, R΄S

This type of reaction is also effective for intramolecular carbon–carbon bond formation. The electrochemical oxidation of α-heteroatom-substituted tetraorganostannanes having a carbon–carbon double bond in an appropriate position gives rise to effective cyclization (Equation 35.39). A fluoride ion derived from the supporting electrolyte (Bu4NBF4 or Bu4NPF6) is introduced to the cyclized carbocation intermediate [153]. The reaction is generally applicable to the formation of six- and seven-membered rings and provides an efficient route to fluorine-containing compounds via carbon–carbon bond formation. SnBu3 Y R n Y = O, NCO2Me

–e– C anode undivided Bu4NBF4 or Bu4NPF6 CH2CI2

+ Y

Y

Y +

R

R n

F

R n

(35.39)

n

n = 1, 2

The bromide ion can also be introduced using cathodically generated Br − as a nucleophile. The use of Br − as the supporting electrolyte results in the selective oxidation of Br − without affecting the substrate. Thus, the anodic oxidation using Bu4NClO4 as a supporting electrolyte in CH2Br2 leads to the formation of the cyclized bromides [154]. Since this reaction does not proceed in a divided cell, the cathodic reduction seems to play an important role. A mechanism involving the reduction of CH2Br2 to generate Br −, which reacts with the anodically generated cyclized cation, has been suggested (Equation 35.40).

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1373

Organoelemental Compounds Anodic process SnBu3

–e– C anode undivided

O

Bu4NCIO4 CH2Br2

R n

+ O

O

+

R

R

(35.40)

n

n

O

Cathodic process CH2Br2

R n

+e–

Br

Br–

The activation by σ–n interaction is also applicable for acyl-elemental compounds such as acylsilanes. The oxidation potentials of acylsilanes are less positive than those of aldehydes and ketones because the nonbonding p orbital of the carbonyl oxygen atom interacts effectively with the C–Si σ orbital [155,156]. This effect seems to be small in the case of aromatic acylsilanes [157]. The anodic oxidation of acylsilanes results in the cleavage of the C–Si bond. Various nucleophiles including water, alcohols, and tosyl amides are introduced to the carbonyl carbon (Equation 35.41) [155,156]. –2e– C anode undivided

O R

SiMe3

NuH Bu4NCIO4 CH2CI2

O R

Nu

(35.41)

NuH = R΄OH, H2O, NHTs

The acylsilane tosylhydrazones exhibit slightly less positive oxidation potentials than the corresponding tosylhydrazones of aldehydes [156]. The preparative anodic oxidation of acylsilane tosylhydrazones in a divided cell gave the corresponding nitriles. This reaction seems to be promoted by the electrogenerated acid. e. Activation by σ–σ Interaction The tetraorganoelemental compounds are also activated toward electron transfer by the interaction of a carbon–element σ orbital with a neighboring carbon–element σ orbital. For example, the oxidation potential of 1,2-bis(trimethylsilyl)ethane is much less positive than those of trimethylsilylethane, 1,1-bis(trimethyl)methane, and 1,3-bis(trimethylsilyl)propane (Table 35.7) [12,158]. The interaction between two neighboring C–Si σ orbitals is responsible for this phenomenon. The geometric requirement of the interaction is important. The oxidation potential of cis(exo,exo)-2,3bis(trimethylsilyl)norbornane exhibits a less positive oxidation potential than the trans isomer, because in the former case, two C–Si bonds are in the same plane, whereas in the latter case, they are in almost perpendicular orientation. The preparative electrochemical oxidation of cis(exo,exo)2,3-bis(trimethylsilyl)norbornane resulted in the cleavage of two C–Si bonds to form norbornene (Equation 35.42). The σ–σ interaction is also effective for organotin compounds (Table 35.7) [12,158]. The oxidation potential of 1,2-bis(tributylstannyl)ethane is much less positive than that of (tributylstannyl)ethane, indicating the significant interaction between two neighboring C–Sn σ orbitals. The interaction between the C–Si σ orbital and the C–Sn σ orbital is, however, less effective

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Organic Electrochemistry

TAbLE 35.7 Oxidation Potentials of 1,2-Disilyl- and 1,2-Distannyl-Substituted Ethanes and Related Compounds Compound

Ed (V vs. Ag/AgCl)

Me3Si Me3Si

Ed (V vs. Ag/AgCl)

Compound

2.19 SiMe3

>2.5

2.20 SiMe3

Me3Si

SiMe3

1.65

1.74 SiMe3

Me3Si

2.02

SiMe3

SiMe3 SiMe3

1.41

1.45

Bu3Sn

Bu3Sn

SiMe3

SnBu3

0.66

1.22

Bu3Sn

Sources: Yoshida, J. and Nishiwaki, K., J. Chem. Soc., Dalton Trans., 2589, 1998; Yoshida, J. et al., Abstract of the 16th Symposium on Electroorganic Chemistry, Tokyo, 1994, p. 83. Ed: Decomposition potentials determined with rotating disk electrode voltammetry in LiClO4/CH3CN using a glassy carbon working electrode.

as demonstrated by 1-trimethylsilyl-2-(tributylstannyl)ethane. This is probably because the energy match between the two orbitals is not good. Si(CH3)3 Si(CH3)3

+

–e– (3.0 F)

Si(CH3)3 Si(CH3)3

CH3OH

(35.42) Si(CH3)3

Si(CH3)3

–e– +

96%

It is interesting that the σ–σ interaction system further interacts with a neighboring π-system [159,160]. For example, the oxidation potential of 1,2-diphenyl-1,2-bis(trimethylsilyl)ethane is less positive than those of benzyltrimethylsilane and 1,2-bis(trimethylsilyl)ethane (see Scheme 35.1). Such interactions seem to provide an important concept for the development of new electronic systems. f. Activation by Dynamic Coordination As a method for the activation of tetraorganoelemental compounds toward electron transfer, the dynamic coordination has received significant research interest because it does not utilize orbital interactions. The utility of this concept is demonstrated by the following example. The oxidation potential of (3-oxobutyl)tributylstannane is less positive than that of tetrabutylstannane (Table 35.8) [161]. Since spectroscopic studies did not indicate the coordination of the carbonyl group in the neutral molecule, the electron transfer must be facilitated by the coordination in the cation radical intermediate. In fact, such dynamic coordination is suggested by ab initio molecular orbital calculations. It is particularly interesting that the effect of the carbonyl group depends on its position. The dynamic coordination to form a five-membered ring seems to be the most effective. The coordination to form a

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1375

Organoelemental Compounds SiMe3

SiMe3

SiMe3

SiMe3

SiMe3 Ed = 1.39V

meso: Ed = 1.23V dl:

Ed = 1.32V

Ed = 1.74V Ed : V vs. Ag/AgCl

SCHEME 35.1 Oxidation potentials of some organosilicon compounds. (From Yoshida, J. and Nishiwaki, K., J. Chem. Soc. Dalton Trans., 2589, 1998.); Ed: Decomposition potentials determined by rotating-disk electrode voltammetry in LiClO4/CH3CN using a glassy carbon working electrode.

six- and seven-membered ring is less effective. The ether group, the pyridyl group, and the phenyl group are also effective as coordinating groups. The effect of the pyridyl group is especially remarkable. The preparative electrochemical oxidation followed by the treatment with aq NaCl results in the cleavage of the Sn–C bond and the introduction of Cl to Sn to give a five coordinated organotin compound (Equation 35.43). In the cation radical intermediate, the C–Sn bond is cleaved to produce the tin cation stabilized by coordination and the alkyl radical. The tin cation is trapped by Cl− upon workup to give tin chloride. The alkyl radical is further oxidized to give the cation, which undergoes an elimination reaction via isomerization to yield a regioisomeric mixture of alkenes. +

–e– C anode divided cell

O SnR3

+

O

O

SnR3

SnR2

+ Bu4NCIO4 CH2CI2

(35.43)

R –e–

CI–

R+

O SnR2

Alkenes

CI

TAbLE 35.8 Oxidation Potentials of Heteroatom-Substituted Tetraorganostannanes Compound

Ed (V vs. Ag/AgCl)

Compound

Ed (V vs. Ag/AgCl)

O 1.41 SnBu3 SnBu3

O

SnBu3

N

SnBu3

1.55

O O

1.42

1.14

1.61 SnBu3 Bu4Sn

1.50 1.67

SnBu3

Source: Yoshida J. and Izawa, M., J. Am. Chem. Soc., 119, 9361, 1997. Ed: Decomposition potentials determined by rotating disk electrode voltammetry in Bu4NClO4/CH2Cl2 using a glassy carbon electrode.

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Organic Electrochemistry

TAbLE 35.9 Oxidation Potentials of the Heteroatom Compounds Containing a 2-Pyridylethyldimethylsilyl group Compound

Ed (V vs. SCE)

Compound

Ed (V vs. SCE)

2-PyeMe2SiCH2OMe

1.36

2-PyeMe 2Si CHOMe  SiMe3

1.19

Me3SiCH2OMe

1.58

Me3Si CHOMe  SiMe3

1.25

2-PyeMe2SiCH2SPh

0.93

2-PyeMe 2Si CHSPh  SiMe3

0.85

Me3SiCH2SPh

1.12

Me3Si CHSPh  SiMe3

1.05

Source: Yoshida, J. et al., Chem. Lett., 251, 1999. Ed: Decomposition potentials determined with rotating disk electrode voltammetry using a glassy carbon anode in LiClO4/CH3CN.

The dynamic coordination is also effective for the activation of α-heteroatom-substituted tetraorganosilanes [162]. The oxidation potentials of the 2-pyridylethyl (2-Pye)-substituted compounds are less positive than those of the corresponding parent compounds (Table 35.9). The decrease in the oxidation potential can be explained in terms of the coordination of the pyridyl group to silicon in the cation radical intermediate. The dynamic coordination also facilitates the selective bond cleavage. In fact, the preparative electrochemical oxidation of a compound having both a (2-Pye)Me2Si group and a Me3Si group resulted in the selective cleavage of the C–SiMe2(2-Pye) bond as shown in the following equation.

SiMe2

N Me3Si

SPh

1.0 V vs. SCE C anode (1.92 F) undivided cell 0.2 M Bu4NBF4 MeOH

+ OMe

N SiMe2 Me3Si

SPh

Me3Si

(35.44)

SPh 90%

2. Electrochemical Oxidation of Dielemental Compounds (R3E–ER3) The group 14 organoelemental compounds having element–element bonds are rather easily oxidized in comparison with simple organoelemental compounds. This can be understood in terms of the high-energy element–element σ orbital [163]. The oxidation potentials of compounds having Si–Si bonds, disilanes, and polysilanes are interesting (Table 35.10). The oxidation potentials decrease with increasing chain length [164]. A decrease in the HOMO level with an increase in the chain length seems to be responsible for this phenomenon. The oxidation potentials of polysilanes also depend on the nature of the organic group on silicon [165]. The preparative electrochemical oxidation of hexamethyldisilane resulted in the cleavage of the Si–Si bond, although the product was not fully characterized [164]. The anodic oxidation of cyclic polysilanes using Et4NBF4 as supporting electrolyte gave rise to ring opening to form α,ω-difluorosilanes (Equation 35.45) [166,167].

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Organoelemental Compounds

TAbLE 35.10 Oxidation Potentials of Permethylpolysilanes Polysilane

Oxidation Potential (V vs. SCE)

Me(SiMe2)2Me Me(SiMe2)3Me Me(SiMe2)4Me Me(SiMe2)5Me Me(SiMe2)6Me

1.88 1.52 1.33 1.18 1.08

Source: Boberski, W.G. and Allred, A.L., J. Organometal. Chem., 88, 65, 1975. Determined by AC polarography.

TAbLE 35.11 Oxidation Potentials of group 14 Diorganoelemental Compounds Compound Me3SiSiMe3 Me3SiGeMe3 Me3GeGeMe3

Ep (V vs. Ag/AgCI)

Compound

Ep (V vs. Ag/AgCI)

1.76 1.76 1.70

Me3SiSnMe3 Me3GeSnMe3 Me3SnSnMe3

1.60 1.44 1.28

Source: Mochida, K. et al., Bull. Chem. Soc. Jpn., 58, 2149, 1985. Ep: Peak potentials determined by cyclic voltammetry in Et4NClO4/CH3CN with a glassy carbon anode.

(SiMe2)n

–e–(4 F) Pt anode divided cell

F–(SiMe2)n–F

Et4NBF4 CH2CI2–CH3CN

(35.45)

n

1

2

3

4

5

6

% Yield

5.5

27

12

49.5

1.5

0.5

The oxidation potentials of other group 14 dielemental compounds are shown in Table 35.11 [168]. The oxidation potentials decrease in the order: Si–Si > Si–Ge > Ge–Ge > Si–Sn > Ge–Sn > Sn–Sn, probably reflecting the relative energy level of the element–element σ orbital. The dynamic coordination facilitates the selective Si–Si bond cleavage as observed in cases of C–Sn and C–Si bonds (Table 35.12) [169]. Oxidation potentials of the disilane having 2-pyridylethyl groups are much lower than those of hexamethyldisilane and tetramethyldiphenyldisilane. The 2-pyridylphenyl group is more effective for decreasing the oxidation potential of disilane because of its decreased flexibility in the conformation. Thus, a significant effect of the 2-pyridyl group on the oxidation potential may be ascribed to effective intramolecular coordination to stabilize the radical cation intermediate as suggested by DFT calculations. 3.

Electrochemical Oxidation of Element–Element Double bonds (R2E═ER2, R2E═ER–ER═ER2) The electrochemical property is one of the most important physical properties of disilene and their derivatives [170]. The oxidation potentials of compounds having Si–Si and Ge–Ge double bonds, disilenes, tetrasila-1,3-butadiene, digermene, and tetragerma-1,3-butadiene are shown in Table  35.13 [170,171].

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Organic Electrochemistry

TAbLE 35.12 Oxidation Potentials of Disilanes Ed (V vs. SCE)

Compound Me3Si – SiMe3

1.38

PhMe2Si–SiMe2Ph

1.33

N Me2Si

SiMe2

1.00

N

Me2Si N

SiMe2

N

0.60

Ed: Decomposition potentials determined by rotating disk electrode voltammetry using a glassy carbon anode in LiClO4/CH3CN. Source: Nokami, T. et. al., Beilstein J. Org. Chem., 3, 7, 2007.

TAbLE 35.13 Oxidation Potentials of Disilene, Tetrasila-1,3-butadiene, and germanium Analogues Ep(ox)(V vs. Ag/AgCl)

Compound R

R 0.56 (E1/2)

Si Si R

R

R

R Si

Si

R

R Si

0.07 (E1/2)

Si

R

R

R

R Ge

0.28

Ge R

R R

R Ge Ge

R

R

0.15

Ge Ge R

R

Determined by cyclic voltammetry in o-dichlorobenzene/Me3PhNB(C6F5)4. R = 2,4,6-triisopropylphenyl. Source: Schäfer, A. et al., Chem. Eur. J., 15, 8424, 2009.

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Organoelemental Compounds

The oxidation potentials decrease in the order: Si=Si > Ge=Ge > Ge=Ge–Ge=Ge > Si=Si–Si=Si. High-lying HOMOs and conjugation between the double bonds in butadienes are suggested. 4. Electrochemical Oxidation of Organoelemental Hydrides (R3EH) Elemental hydrides are susceptible to oxidation in the absence of activation. For example, the electrochemical oxidation of PhMe2SiH in the presence of CuCl leads to the formation of PhMe2SiCl [172]. The use of R2SiH2 in the absence of CuCl results in the formation of oligosilanes as shown in the following equation [173,174]. –e– Pt electrode undivided cell

R1

R1R2SiH2

H

(35.46)

H

Si

Bu4NBF4 DME

R2 n

The anodic oxidation of tin hydrides has also been studied. The anodic oxidation of Ph3SnH leads to the formation of the corresponding tin radical (+0.80 V vs. SCE) [175,176]. The tin radical dimerizes to form Ph3SnSnPh3. The tin radical can also be utilized for the initiation of radical chain reactions promoted by tin hydride as demonstrated by Equation 35.47. At more positive potential (+1.15 V), Ph3SnH is reported to be oxidized to tin cation. Br

Ph3SnBr

O

–e– Ph3SnH

O

(35.47)

Ph3Sn

O

O Ph3SnH

. 5. Electrochemical Oxidation of Stable Radicals of the Heavy group 14 Elements (R3E ) Recently, stable free radicals of the heavy group 14 elements became accessible by protecting the radical center with large organosilicon substituents. The oxidation potentials of these isolable radicals are shown in Table 35.14 [177]. Although all of the first oxidation potentials are irreversible in both o-dichlorobenzene (o-DCB) and THF, the oxidation potentials decrease in the order: Si > Ge > Sn, which is consistent with experimental and computed ionization energies. TAbLE 35.14 Oxidation Potentials of Stable Radicals of the Heavy group 14 Elements Radicals ( Bu2MeSi)3Si. (tBu2MeSi)3Ge. (tBu2MeSi)3Sn. t

CV measurements were performed in o-DCB/Bu4NB(C6F5)4. Oxidation potentials in parentheses were measured in THF/Bu4NBPh4. Source: Becker, J.Y. et al., Chem. Eur. J., 15, 8480, 2009.

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Ep (ox) (V vs. Ag/AgCl) 0.40 (0.8) 0.28 (0.6) −0.05 (0.2)

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Organic Electrochemistry

6. Electrochemical Oxidation of Silylenes and germylenes ((R2N)2E:) The oxidation potentials of silylenes and germylenes are shown in Table 35.15 [178]. The oxidation potentials decrease in the order: saturated germylenes > unsaturated germylenes > saturated silylenes > unsaturated silylenes. Unsaturated germylenes and silylenes, which have a 6π electron system, are oxidized at lower oxidation potentials than those of saturated analogues. However, the measured oxidation potentials do not correlate with the ionization potentials because of kinetic and surface effects in the solution electrochemistry. 7. Electrochemical Oxidation of Other Types of group 14 Organoelemental Compounds This section is concerned with cases where the initial electrochemical reaction does occur at a functional group other than the organoelemental part in the substrate, and the subsequent chemical or electrochemical process takes place at the organoelemental part. In other words, the organoelemental part plays a major role in follow-up processes. The first example of this category is the anodic oxidation of 1-trimethylsilyl-1,3-dienes [179]. The initial oxidation of the 1,3-diene part in methanol generates the 1,4-dimethoxylated intermediate just as simple 1,3-dienes (Equation 35.48). The electronic effect of the α-silyl group to the π-system is not so large as far as the oxidation is concerned, although the α-silyl group slightly lowers the LUMO of the π-system (see Section IV). The 1,4-dimethoxylated intermediate can also be seen as an allylsilane, and therefore undergoes the electrochemical oxidation in a similar fashion to simple allylsilanes. The C–Si bond is cleaved, and a methoxy group is introduced on the carbon. In this case, the methoxy group is introduced selectively at one of the allyl terminals bearing the methoxy group that is introduced by the first oxidation. This selectivity is explained in terms of the cation-stabilizing effect of the methoxy group. Consequently, the overall reaction leads to the formation of 1,1,4-trimethoxy-2-butene derivatives, and this reaction provides a useful transformation in organic synthesis. –2e– C anode undivided cell Ph

Et4NOTs MeOH

SiMe3

OMe SiMe3

Ph

OMe

–2e–

OMe

Ph OMe

OMe

(35.48) TAbLE 35.15 Oxidation Potentials of Silylenes and germylenes Ep Silylene (V vs. Ag/AgCl)

Ep Germylene (V vs. Ag/AgCl) tBu

tBu

N

N Si

0.67

Ge

N

N

tBu

tBu

tBu

tBu

N

0.84

N Si

N tBu

0.95

Ge

1.20

N tBu

CV measurements were performed in THF/Bu4NClO4. Source: Dhiman, A. et. al., Organometallics, 23, 5689, 2004.

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Organoelemental Compounds

Another example of the anodic oxidation of π-systems having a silyl group at the α position is the electrochemical oxidation of 2,5-bis(trimethylsilyl)thiophene resulting in the elimination of the trimethylsilyl group to produce polythiophene. Polythiophenes obtained by this method have high conductivity [180]. The anodic oxidation of β-silylcarboxylic acids is fascinating. The reaction results in the elimination of the carboxylate group and the silyl group to form olefinic products (Equation 35.49) [117,181]. A plausible mechanism involves the initial electrochemical oxidation to eliminate CO2 followed by facile β-elimination of the silyl group from the thus generated carbocation to form a carbon–carbon double bond. –2e– C anode

R SiMe3

R

CH3CN/MeOH (5:1)

CO2Na

R +

(35.49)

SiMe3

The anodic oxidation of α-silylacetic acids at a Pt anode leads to the formation of Kolbe dimers (Equation 35.50) [182]. This reaction serves as a powerful method for synthesizing 1,2-disilylethanes [183]. The electrolysis using a carbon anode leads to two competing reactions: decarboxylation and desilylation [184].

R3Si

CO2H

–e– Pt anode

R3Si

(35.50)

SiR3

CH3CN/MeOH (3:1)

Facile β-elimination of the silyl group is also utilized in intramolecular anodic olefin coupling reactions [185–187]. For example, the intramolecular anodic coupling of an enol ether with an allylsilane group has been reported (Equation 35.51). This reaction seems to be quite useful for the construction of functionalized cyclic compounds because it leads to the regioselective formation of olefinic products via facile β-silyl elimination. R΄

MeO R –e– SiMe3

n R΄

LiCIO4 MeOH 2,6–lutidine

MeO

R

MeO

n

(35.51)

H

IV. ELECTROCHEMICAL REDUCTION OF ORgANOELEMENTAL COMPOUNDS In general, it is rather difficult to reduce main group organoelemental compounds, especially group 1 and 2 organoelemental compounds. Group 13 elemental compounds, however, are reducible. Group 14 elemental halides are also electrochemically reducible under suitable conditions, like organic halides. Another interesting aspect of electrochemical reduction of organoelemental compounds is that group 14 elemental groups such as silyl groups behave as electron-withdrawing group to neighboring π-systems by virtue of the so-called dπ−pπ interaction. Such systems are susceptible toward electrochemical reduction, and elements play roles of both activating and controlling groups. Therefore, this section mainly focuses on the electrochemical reduction of group 14 organoelemental compounds.

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A.

Organic Electrochemistry

ELECTROCHEMICAL REDUCTION OF GROUP 13 ORGANOELEMENTAL COMPOUNDS

Triphenylborane exhibits a one-electron reduction wave (E1/2= −2.61 V vs. Ag/Ag+) corresponding to the formation of the anion radical species at a dropping mercury electrode in THF [188]. A second polarographic wave corresponding to the dianion occurs at E1/2= −3.67 V vs. Ag/Ag+. Cyclic voltammetric studies indicate that the anion radical formation is quasireversible, whereas the dianion formation is irreversible. The carbon–boron bonds can be cleaved by cathodic reduction. Cathodic reduction of arylboronic acids and esters under an oxygen atmosphere gives the corresponding phenol derivatives (Equation 35.52) [189]. A mechanism involving cathodically generated superoxide ions has been suggested. 1) +e–, O2 Bu4NCIO4 CH3CN divided cell MeS

B(OH)2

MeS

(35.52)

OH

2) H3O+ 91%

B.

ELECTROCHEMICAL REDUCTION OF GROUP 14 ORGANOELEMENT HALIDES

1.

Reduction Potentials and general Reaction Patterns of group 14 Organoelement Halides It is rather difficult to determine the reduction potentials of halosilanes because they are very easily hydrolyzed. Hydrogen halides produced by the hydrolysis exhibit less negative reduction potentials, and this may cause some difficulty in the measurements [190,191]. Careful measurements in anhydrous solvents, however, revealed that reduction potentials of halosilanes are generally more negative than the corresponding organic halides (Table 35.16) [192,193]. The reduction potentials of other group 14 elemental halides have also been determined [192,193]. As to the nature of the elements, it is notable that the reduction becomes more facile in the order of Si < Ge < Sn < Pb [194]. It should also be kept in mind that the reduction potential becomes less negative as one progresses the series F, Cl, and Br. The general reaction pattern for the reduction of group 14 organoelement halides is shown in Equation 35.53 [194]. The one-electron reduction of elemental halides leads to the formation of element-centered radicals. The element radicals may undergo radical coupling to produce metal– metal bonds or addition to unsaturated compounds. The element radical may also be reduced electrochemically under some conditions to give the anions, which react with electrophiles.

TAbLE 35.16 Polarographic Reduction Potentials of Halosilanes and Halogermanes Compound Ph3SiF Ph3SiCl Ph3SiBr

E1/2 (V)

Compound

E1/2 (V)

−2.15 −1.95 −1.93

Ph3GeF Ph3GeCl Ph3GeBr

−1.85 −1.75 −1.60

Source: Corriu, R.J.P. et al., J. Organometal. Chem., 188, 63, 1980. E1/2 values were measured vs. Bu4NI saturated AgI/Ag reference electrode.

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Organoelemental Compounds

E

X

+e–

+e–

E

Nucleophilic reaction

E–

X–

(35.53) Addition to unsaturated compounds

Radical coupling

2. Reductive Formation of Element–Element bonds The electrochemical reduction of group 14 elemental halides to form element–element bonds is rather common. The reactions proceed either via radical coupling or via nucleophile attack of element anions to elemental halides, depending on the reaction conditions such as cathode potentials. The electrochemical reduction of chlorosilanes has received significant research interests, because this reaction provides an efficient route to disilanes and polysilanes having Si–Si bonds (Equation 35.54) [190,194–204]. The electrode material and the supporting electrolyte seem to be important factors for the effective homocoupling. Sacrificial anodes have found to be quite effective for the reductive coupling of chlorosilanes, and extensive work has been devoted to this field. In fact, various sacrificial anodes such as Hg [195,199], Mg [200], Cu [201,202], Ag [199], and Al [202–204] were found to be effective for the reductive coupling of chlorosilanes to form disilanes. The application of the method to dichlorosilanes gives rise to the formation of polysilanes (Equation 35.55). R1 2 R2

+e– Si Cl R3

R1 n Cl

R2

R1

R1

Si

Si

R3

R3

(35.54)

R2

R1 +e–

Si

(35.55)

Si

Cl

R2

R2 n

Copolymerization of two different dichlorosilanes can also be achieved by the reduction using a sacrificial anode [205,206]. It is interesting that copolymerization of sila-functional dichlorosilanes with simple dichlorosilanes took place smoothly to provide an elegant route to functionalized polysilanes [206]. As to the synthesis of sila-functional polysilanes, the electrochemical reduction of perfluoroalkyl-substituted trichlorosilane to give perfluoroalkyl-substituted polysilane is interesting [207]. Sequence-ordered polysilanes were prepared by the electrochemical reduction of dichloro-oligosilanes using an Mg sacrificial anode [208]. The reductive coupling of monomers having two chlorosilyl groups leads to the formation of polymers containing Si–Si bonds as shown in the following scheme [209,210]. R1 Cl

R1 +2e–

Si R2

CH2

CH2

Si R2

Cl

CH2

CH2

R1

R1

Si

Si

R2

R2

(35.56) n

The Mg sacrificial anode is also effective for the synthesis of Si–Ge and Ge–Ge bonds. This method is quite successful for the preparation of polygermanes and germane–silane copolymers (Equation 35.57) [211]. Polygermanes were also prepared by conventional electrochemical reduction of dihalogermanes using a Pt cathode and a Ag anode [212].

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1384

Organic Electrochemistry R1 Cl

R1

Si

Cl

+

R2

Cl

Ge Cl

R1

R1

Si

Ge

R2

2

+2e–

R2

m

R

(35.57) n

The use of sacrificial anodes inevitably leads to the formation of elemental halides as by-products. This problem can be solved by the use of a silicon carbide electrode. The homocoupling took place smoothly in an undivided cell without the formation of organoelement halides [213]. A hydrogen electrode was also effective for this purpose [213]. Under suitable conditions, the electrochemical reduction of R2SiCl2 resulted in the formation of a Si–Si double bond rather than polymerization. The cathodic reduction of dichlorodimesitylsilane with a Hg pool cathode and a Pt anode in DME (dimethoxyethane) gave tetramesityldisilene (Equation 35.58) [214]. The bulky substituents on the silicon atom seem to be responsible for the retardation of the polymerization and the protection of the highly reactive Si–Si double bond.

Cl Si

Cl

+e– Hg cathode Ag anode Si

Si

Bu4NClO4 DME divided cell

(35.58)

20%

Cathodic reduction is also useful for the formation of Sn–Sn bonds [194]. For example, Ph3SnCl exhibits two major reduction peaks (Hg cathode, −1.40 and −2.60 V vs. SCE in cyclic voltammetry [215,216]). At lower potential, Ph3SnCl undergoes a one-electron reduction to give the tin radical that dimerizes to give Ph3SnSnPh3. At higher potentials, a two-electron reduction process takes place to give the tin anion that reacts with Ph3SnCl to produce Ph3SnSnPh3. Some modifications have been reported for the electrochemical dimerization of R3SnCl [80,217]. In a similar fashion, tin formates (R3SnOCHO) also undergo a cathodic reduction to give the corresponding homocoupling products as shown in the following equation [218].

Ph3SnOCHO

+2e–(2.5 F) Stainless steel cathode undivided cell Dimethylacetamide

(35.59)

Ph3SnSnPh3 98%

The electrochemical reduction of a mixture of a R3SnCl and a R′3SiCl using Mg electrodes under ultrasonic irradiation was reported to give the corresponding cross-coupling product together with the distannane (Equation 35.60) [219]. The disilane was not formed. The fact that the reduction potential of R3SnCl is less negative than that of R′3SiCl implies that the reaction proceeds by the initial reduction of the R3SnCl. The product distribution is also consistent with this mechanism.

Bu3SnCl

+ Me3SiCl 10 equiv

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+e– Mg cathode Mg anode Undivided cell LiClO4 THF

Bu3SnSiMe3 50%

+ Bu3SnSnBu3 41%

+ Me3SiSiMe3 0%

(35.60)

1385

Organoelemental Compounds

3. Reductive Formation of Element–Carbon bonds The reduction of group 14 elemental halides in the presence of organic compounds of more negative reduction potentials leads to the formation of the element–carbon bond. This type of reaction also provides a method for the synthesis of organoelemental compounds, although there are only a few examples. A mechanism involving the initial formation of the element anion intermediate followed by the nucleophilic reaction with organic substrates is proposed. Another possibility to be considered is the formation of the element-centered radical intermediate, which reacts with organic substrates. The electrochemical reduction of R2SiCl2 in the presence of 2,3-dimethylbutadiene proceeds smoothly to give the sila-cyclopentene derivatives (Equation 35.61) [220]. Probably, the initial reduction of R2SiCl2 produces the silyl anion, which adds to the 1,3-diene. The intramolecular displacement reaction of the resulting allylic carbanion with a chlorosilane moiety gives the cyclized product.

Ph2SiCl2

+2e–(2.8 F) Pt cathode Cu anode

+

Bu4NBPh4 THF

Ph Ph

+

Si

Si Ph

Cl

Polymer 44%

Ph

28%

(35.61) The following reaction also provides an example of the reactions of cathodically generated silyl species with organic compounds. The electrochemical reduction of R3SiCl in the presence of unsaturated compounds such as phenylacetylene, styrene, and cyclohexene gives the corresponding formal hydrosilylation products (Equation 35.62) [221,222]. In this case, a mechanism involving the silyl radical seems to be more plausible because the addition of the silyl anion to cyclohexene is quite unlikely. The hydrogen that was introduced to another formal acetylenic or olefinic carbon seems to come from the solvent via radical abstraction. The disilane is also effective as a source of the silyl group. +2e–(2.3 F) Pt cathode Me3SiCl

+

Me3Si

(35.62)

Divided cell 72%

The tin anion generated by cathodic reduction can also be utilized for carbon–element bond formation. The electrochemical reduction of mixtures of R3SnCl in the presence of organic chlorides gives the corresponding coupling products (Equation 35.63). Since the reduction potentials of R3SnCl are less negative than those of organic chlorides, the reaction seems to proceed by the initial reduction of R3SnCl to generate the stannyl anion, which attacks organic chlorides nucleophilically to give the final product [219].

Bu3SnCl

+

+2e– Mg cathode Mg anode BuCl 5 equiv

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Undivided cell LiClO4/THF

Bu3Sn–Bu 76%

+

Bu3SnSnBu3 20%

+ Bu–Bu Trace

(35.63)

1386

C.

Organic Electrochemistry

ELECTROCHEMICAL REDUCTION OF GROUP 14 ORGANOELEMENTAL COMPOUNDS CONTAINING REDUCIBLE π- AND σ-SYSTEMS

It is known that group 14 elemental compounds facilitate the electron transfer from the neighboring π-system by virtue of the so-called dπ–pπ interaction [223]. The LUMO level is decreased by such an interaction, which in turn favors the reduction of the system. The silyl group also stabilizes the anion radical intermediate formed by the one-electron reduction [224,225]. These effects can also be explained in terms of the orbital interaction with the low-lying σ* orbital of the silyl group instead of the vacant d orbital. Arylsilanes serve as a typical example of this system. The reduction potentials of arylsilanes are slightly less negative than those of the parent aromatic hydrocarbons [226–230]. This seems to be attributed to the dπ–pπ interaction between the aromatic ring and the silicon atom. The electrochemical behavior of silyl-substituted cyclooctatetraene is interesting [231]. The second reduction potential becomes less negative by the silyl-substitution. The stabilization of the dianion (aromatic 10 π-system) by dπ–pπ interaction seems to be responsible for this phenomenon. Preparative electrochemical reduction of arylsilanes in methylamine gives Birch-type products (Equation 35.64) [225]. The hydrogen atom is introduced on the carbon adjacent to the silyl group preferentially, and this regioselectivity is explained in terms of the stabilization of the anion radical by the neighboring silyl group.

SiMe3

+2e– Pt cathode undivided cell LiCl CH3NH2

SiMe3

SiMe3 +

94%

(35.64)

3%

Reduction of α-halo organoelemental compounds (R3MCH2X) also seems to be assisted by a similar orbital interaction. The negative charge formed by the reduction of the carbon–halogen bond should be stabilized by the interaction with group 14 elements. As a matter of fact, the reduction potentials of halomethylsilanes are less negative than those of simple alkyl halides [232,233]. The electrochemical reduction of halomethylstannanes has also been reported [234].

D.

ELECTROCHEMICAL REDUCTION OF GROUP 15 ORGANOELEMENTAL COMPOUNDS

As other examples of electrochemical reduction of organoelemental compounds, the reduction of organophoshines such as triphenylphosphine and triphenylphosphine oxide has been investigated. The polarographic reduction of triphenylphosphine proceeded by a one-electron transfer to form the anion radical that undergoes cleavage of a phenyl group as shown in the following equation [235].

Ph3P

Polarographic reduction

Ph2PH + Ph – Ph

(35.65)

In a similar way, the electrochemical reduction of triphenylphosphine oxide led to the cleavage of the P–C bond, and a complex mixture of diphenylphosphine oxide, diphenylphosphine, phenylphosphine oxide, etc.; however, electroreduction of triphenylphosphine dichloride in acetonitrile was performed successfully in an undivided cell fitted with an aluminum sacrificial anode to give triphenylphosphine [236,237]. The one-pot transformation of triphenylphosphine oxide to triphenylphosphine was also achieved successfully by the treatment of triphenylphosphine oxide in acetonitrile with oxalyl chloride and subsequent electrochemical reduction (Equation 35.66).

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1387

Organoelemental Compounds

Ph3P=O

(COCl)2

Cl Ph3P

Cl

+e– Al cathode undivided cell AlCl3 CH3CN

Ph3P

(35.66)

74%

Alternatively, the electroreduction of triphenylphosphine oxide to triphenylphosphine in an acetonitrile solution in the presence of chlorotrimethylsilane was performed successfully in an undivided cell using a zinc sacrificial anode (Equation 35.67) [238,239]. These results would offer facile and environmentally benign methods for the recycling of triphenylphosphine oxide, which is inevitably produced in synthetic processes such as the Wittig reaction and Mitsunobu reactions. +2e– Zn cathode undivided cell Ph3P=O

Bu4NBr CH3CN

Ph3P

(35.67)

89%

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Organoelemental Compounds 183. 184. 185. 186. 187. 188. 189. 190. 191. 192. 193. 194. 195. 196. 197. 198. 199. 200. 201. 202. 203. 204. 205. 206. 207. 208. 209. 210. 211. 212. 213. 214. 215. 216. 217. 218. 219. 220. 221. 222. 223. 224. 225. 226. 227. 228. 229. 230. 231. 232.

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233. Mairanovskii, S. G.; Ponomarenko, V. A.; Barashkova, N. V.; Kaina, M. A. Izv. Akad. Nauk SSSR, Ser. Khim. 1964, 1951–1956; Chem. Abstr. 1965, 62, 6378c. 234. Devaud, M.; Lecat, J.-L. Bull. Soc. Chim. Fr. 1985, 1187–1190. 235. Santhanam, K. S. V.; Bard, A. J. J. Am. Chem. Soc. 1986, 90, 1118–1122. 236. Yano, T.; Kuroboshi, M.; Tanaka, H. Tetrahedron Lett. 2010, 51, 698–701. 237. Kuroboshi, M.; Yano, T.; Kamenoue, S.; Kawakubo, H.; Tanaka, H. Tetrahedron 2011, 67, 5825–5831. 238. Tanaka, H.; Yano, T.; Kobayashi, K.; Kamenaue, S.; Kuroboshi, M.; Kawakubo, H. Synlett 2011, 582–584. 239. Kawakubo, H.; Kuroboshi, M.; Yano, T.; Kobayashi, K.; Kamenoue, S.; Akagi, T.; Tanaka, H. Synthesis 2011, 4091–4098.

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36

Organometallic Compounds as Tools in Organic Electrosynthesis Anny Jutand

CONTENTS I. II.

III.

IV.

Introduction ........................................................................................................................ 1394 One-Electron Activation of Organometallic Complexes in Transition Metal–Catalyzed Electrosyntheses ................................................................................................................. 1395 A. Electroreductive Carboxylation of Aryl Halides Catalyzed by Nickel ....................... 1395 B. Electroreductive Carboxylation of α-Methylbenzyl Chlorides Catalyzed by Nickel .... 1397 C. Electroreductive Homocoupling of Aryl/Vinyl Halides Catalyzed by Nickel ........... 1398 1. Electroreductive Homocoupling of Aryl Halides Catalyzed by Nickel............... 1398 2. Electroreductive Homocoupling of Vinyl Halides Catalyzed by Nickel ............. 1400 D. Electroreductive Polymerization of Organic Dihalides Catalyzed by Nickel ............ 1401 E. Electroreductive Heterocoupling of Two Organic Halides Catalyzed by Nickel ....... 1402 F. Electroreductive Heterocoupling of Aryl Halides with Vinylic/Allylic Acetates or Aryl Halides Catalyzed by Cobalt.......................................................................... 1404 Two-Electron Activation of Organometallic Complexes in Transition Metal–Catalyzed Electrosyntheses ................................................................................................................. 1406 A. Electroreductive Carboxylation of Aryl Halides/Triflates Catalyzed by Palladium .......1406 1. Electroreductive Carboxylation of Aryl Halides Catalyzed by Palladium .......... 1406 2. Electroreductive Carboxylation of Aryl Triflates Catalyzed by Palladium ......... 1407 B. Electroreductive Carboxylation of Vinyl Triflates Catalyzed by Palladium ...............1410 C. Electroreductive Cleavage of Allylic Esters Catalyzed by Palladium or Nickel .........1412 D. Electroreductive Homocoupling of Aryl Halides/Triflates Catalyzed by Palladium ............................................................................................................1412 1. Electroreductive Homocoupling of Aryl Halides Catalyzed by Palladium..........1412 2. Electroreductive Homocoupling of Aryl Triflates Catalyzed by Palladium.........1413 E. Electroreductive Heterocoupling of Two Aryl Halides Catalyzed by Palladium ........1414 F. Electrosynthesis of Ketones via the Electroreductive Heterocoupling of Organic Halides Catalyzed by Nickel .......................................................................................1415 G. Electrosyntheses of Cyclopropanes Catalyzed by Nickel or Copper ...........................1416 H. Electroreductive Hydroalkylation of Electron-Deficient Alkenes Catalyzed by Cobalt ..................................................................................................................1417 Electrochemical Recycling of the Catalyst in Transition Metal–Catalyzed Organic Electrosyntheses ..................................................................................................................1418 A. Direct Electrochemical Recycling of the Catalyst .......................................................1418 1. Electroreductive Hydrocarboxylation of Alkynes Catalyzed by Nickel ...............1418 2. Electroreductive Hydroarylation of Alkenes Catalyzed by Nickel or Cobalt .......1419 3. Electrooxidative Carbonylation of Amines Catalyzed by Palladium .................. 1420

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B.

Mediated Electrochemical Recycling of the Catalyst................................................. 1421 1. Electrooxidation of 1,3-Dienes Catalyzed by Palladium ..................................... 1421 2. Electrooxidative Heck-Type Reactions from Arenes Catalyzed by Palladium .....1422 3. Electrooxidation of Alcohols Catalyzed by Palladium ........................................ 1423 4. Electrooxidative Homocoupling of Arylboron Derivatives Catalyzed by Palladium ............................................................................................................. 1424 5. Electrooxidative Wacker Reactions Catalyzed by Palladium .............................. 1425 6. Electrooxidative Asymmetric Dihydroxylation of Alkenes Catalyzed by Osmium... 1426 7. Electroreductive Carbonylation of Aryl Halides Catalyzed by Palladium .......... 1427 8. Palladium-Catalyzed Homocoupling of Aryl Halides Mediated by Electrogenerated Quinoid C8V0 ........................................................................... 1428 V. Conclusion .......................................................................................................................... 1428 References .................................................................................................................................... 1429

I. INTRODUCTION Transition metals were introduced in the late 1960s to catalyze organic reactions in homogeneous processes [1–3]. The transition metal catalyst not only accelerates the reaction but allows (by means of its ligands) a fine control of the chemo-, regio-, and enantioselectivity of the catalytic reaction [1–3]. The mechanism of such catalytic reactions that do not involve any electron transfer may be nevertheless investigated by means of electrochemical techniques, as illustrated by Amatore and Jutand [4–7]. Indeed, most organometallic species involved in a catalytic cycle are electroactive and can be detected and characterized by their reduction (or oxidation) potential. Moreover, their reactivity in elementary steps can be followed by means of electrochemical techniques, since reduction (or oxidation) currents are proportional to their concentration [8–10]. The association of electrochemical techniques (which provide kinetic data and information on the oxidation state of organometallic species) with other techniques such as NMR spectroscopy and ESI MS (which provide structural information) can be used to determine the mechanism of a catalytic reaction involving electroactive organometallic catalysts even if the reaction does not require any electron transfer [4–7]. Transition metals were introduced later on in the early 1980s to catalyze organic electrosyntheses in which electrons are involved [7,11–16]. Transition metal catalysts are required when the organic substrates are not electroactive or when their direct activation by electron transfer does not generate the desired product due to a nondesired bond cleavage. A first activation of an organic substrate by a transition metal complex leads to an intermediate organometallic species which can be activated by electron transfer (one or two electrons). The catalytic cycle is thus a succession or alternation of elementary steps involving chemical activation of organic substrates by transition metals and activation of organometallic intermediates by electron transfer. Those two kinds of activation, chemical activation of organic substrates by a transition metal followed by activation by electron transfer (1 or 2e) of intermediate organometallic species formed in chemical activation steps, are often required to achieve the desired electrosynthesis. Some transition metal–catalyzed organic electrosyntheses do not require any activation of intermediate organometallic species by electron transfer, but electrons may be required to recycle the active catalyst (which initiates the catalytic cycle) from an unreactive one generated in every catalytic cycle, leading to a stoichiometric consumption of electrons. The recycling of the catalyst may be either a direct or a mediated electrochemical process. A transition metal mediator may be used in organic electrosyntheses whose reduction (or oxidation) generates an active species able to reduce (or oxidize) the substrate by an outer-sphere mechanism without any formation of intermediate organometallic species [9,11,12,15–17]. Those homogeneous redox processes give back the initial form of the mediator, which is subsequently transformed into its active form at a cathode (or anode). The electrons are not transferred directly

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from the electrode to the substrate but via the mediator used in catalytic amount. The mediated organic electrosyntheses consume a stoichiometric amount of electrons. Even if the mediator may be a transition metal salt or complex, those catalytic reactions do not involve any intermediate organometallic species, this is why such reactions will not be considered herein. As in any electrosynthesis, two simultaneous electrochemical reactions take place in a transition metal–catalyzed electrosynthesis, one at the cathode and the other at the anode. They can be carried out in a one-compartment cell (noted: cathode, anode in the schemes) or in a two-compartment cell (divided cell) separated by a frit or a membrane (noted: cathode//anode in the schemes). A  one-compartment cell may be equipped with a sacrificial metallic anode whose anodic dissolution releases cations that may have a beneficial role on the catalytic reaction performed at the cathode. An electrosynthesis catalyzed by a transition metal may be performed at constant current (galvanostatic mode) or at controlled potential. In that case, preliminary electrochemical experiments are required to determine the potential imposed during the electrolysis. Transition metal–catalyzed organic electrosyntheses are reported herein. Their mechanisms when known are included. They have been investigated by means of electrochemical techniques (fast cyclic voltammetry and double potential step chronoamperometry at a disk microelectrode, voltammetry and chronamperometry at a rotating disk electrode (RDE) or at an ultramicroelectrode [8–10]). The electrochemical steps involving organometallic complexes have been characterized by the potential and by the number of electron(s) involved in the electron transfer [4]. The organometallic complexes involved in chemical steps have been characterized by their reduction or oxidation potential and their reactivity monitored by the evolution with time of their reduction or oxidation current, leading to the determination of the rate constants.

II.

ONE-ELECTRON ACTIVATION OF ORgANOMETALLIC COMPLEXES IN TRANSITION METAL–CATALyZED ELECTROSyNTHESES

A.

ELECTROREDUCTIVE CARBOXYLATION OF ARYL HALIDES CATALYzED BY NICKEL

The bielectronic reduction of aryl halides formally generates the aryl anions, which could react with carbon dioxide to give aromatic carboxylic acids, as aryl Grignard reagents do (Scheme 36.1). However, such electroreductive carboxylations performed at a Hg pool cathode proceed with poor yields [18]. Moreover, the electrocarboxylation of most aryl bromides or chlorides is problematic since they exhibit very negative reduction potentials in usual organic solvents, acetonitrile, tetrahydrofuran (THF), or dimethylformamide (DMF). Their reduction potential may be even more negative than the reduction potential of CO2 whose reduction becomes the major process [18]. This problem has been partly solved by the introduction of electrogenerated cations (Al3+, Zn2+, Mg2+) that facilitate the electrochemical reduction of aryl halides by complexation of the halides. The cations are supplied by the oxidation of a sacrificial metallic anode (M → Mn+ + ne) in an undivided cell, as pioneered by Silvestri et al. [19] and Périchon et al. [20]. A more versatile alternative is the use of a transition metal complex that catalyzes the electrocarboxylation of aryl halides, such as NiCl2(PPh3)2 [21] or more efficiently NiCl2(dppe) (dppe: 1,2-diphenylphosphinoethane, PPh2–(CH2)2–PPh2), a NiII complex ligated by a bisphosphine ligand

ArCO2– + X–

ArX + CO2 + 2e X = l, Br

SCHEME 36.1

2e X– + Ar–

CO2

Electroreductive carboxylation of aryl halides.

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ArBr + CO2 + 2e 2–20 mmol 1 atm

NiCl2(dppe) 5–10 mol%

ArCO2– + Br– 57–84%

THF + HMPA, rt cathode C // Li anode

Ar = p-Z–C6H4 (Z = C6H5O, CF3, F, H)

SCHEME 36.2 Nickel-catalyzed electroreductive carboxylation of aryl bromides.

(Scheme 36.2) [22], as reported by Fauvarque et al. The electrolyses are performed at less negative potentials than the reduction potential of the aryl halides. The reactions are regiospecific. The mechanism of the catalytic reaction has been elucidated by Amatore and Jutand by means of electrochemical techniques (Scheme 36.3) [23,24]. Most steps of the catalytic cycle have been characterized including the determination of their rate constant. The unreactive precursor NiIICl2(dppe) is reduced in two successive one-electron steps (R1 then R2) at the very beginning of the electrolysis affording a low-ligated complex Ni0(dppe). The latter undergoes a fast oxidative addition to ArBr to give ArNiIIBr(dppe) (Scheme 36.3). The rate constant of this fast oxidative addition has been determined by performing fast cyclic voltammetry at a disk electrode (k1oa = 1.1 × 105 M−1s−1 for PhBr at 20°C). The reduction of ArNiIIBr(dppe) takes place at a less negative potential (R3) than the reduction potential of the corresponding ArBr (e.g., Ep = −1.73 V for PhNiBr(dppe) and −2.7 V vs. saturated calomel electrode [SCE] for PhBr in THF-HMPA [hexamethylphosphoric acid triamide]). The reduction involves one electron and generates a nickel(I) complex: ArNiI(dppe) that reacts with CO2 (k2oa) leading to ArNiIII(μ2-CO2)(dppe) (Scheme 36.3). A reductive elimination (k3re) from this complex affords a NiI carboxylate ArCO2NiI(dppe). The electrolyses are performed at the reduction potential of ArNiIIBr(dppe) (R3). At that potential, ArCO2NiI(dppe) is reduced back to the active Ni0(dppe) by a one-electron transfer and releases the desired aromatic carboxylate. NiIICl2(dppe) –Cl–

1e

THF + HMPA, 20°C L2 = dppe

(R1)

ArX = PhBr

NiICIL2 –Cl–

1e

5 –1 –1 koa 1 = 1.1 × 10 M s

(R2)

ArCO2–

5 –1 –1 koa 2 = 1.5 × 10 M s –1 kre 3 = 100 s

ArX

Ni0L2 1e koa 1 NiIL

ArCO2 Reductive elimination

Oxidative addition

2

kre 3 ArNiIIXL2

Ar–NiIIIL

2

C

O

O

1e

koa 2

(R3)

“Oxidative addition” X– CO2

SCHEME 36.3

ArNiIL2

Mechanism of the nickel-catalyzed electroreductive carboxylation of aryl halides.

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As a consequence of the one-electron transfer, the carboxylation takes place within the coordination sphere of the nickel and not from an aryl anion Ar− (Scheme 36.3) [23,24]. The catalysis is due to the fact that ArNiIIBr(dppe) is more easily reduced than ArBr and CO2 [23,24]. The evolution of the catalytic reduction plateau current of PhNiIIBr(dppe) (measured at an RDE) vs. ArBr and CO2 concentrations allows the determination of the rate constants k2oa and k3re (Scheme 36.3). By comparing the values of k1oa [PhBr], k2oa [CO2], and k3re determined for PhBr, it emerges that the reductive elimination is rate determining for the concentrations of PhBr and CO2 used in the catalytic reactions (≫10 −3 M) [23,24]. The electrolyses consume two electrons per mole of ArBr, one for the monoelectronic activation of ArNiIIBr(dppe) and the other for the recycling of the Ni0 catalyst from ArCO2NiI(dppe). Importantly, electrons are also required for the reduction of the unreactive precursor NiIICl2(dppe) at the very beginning of the electrolysis. A transient and very reactive low-ligated Ni0(dppe) is generated. Its reactivity in oxidative addition to aryl halides has been investigated by performing fast cyclic voltammetry at a disk electrode, leading to the determination of the rate constant (vide supra). This is the only technique that allows both the generation and the characterization of the reactivity of a transient and very reactive low-ligated transition metal complex that cannot be stored.

B.

ELECTROREDUCTIVE CARBOXYLATION OF α-METHYLBENzYL CHLORIDES CATALYzED BY NICKEL

Aryl-2 propionic acids (fenoprofen, ibuprofen, naproxen) are commonly prescribed anti-inflammatory agents (Scheme 36.4). The electrochemical reduction of ArCH(CH3)Cl in the presence of CO2 affords aryl-2 propionic acids, via the carboxylation of the electrogenerated anion ArCH(CH3)−. However, the reactions take place at quite negative reduction potentials [25] but are facilitated by the presence of cations (Zn2+, Mg2+) generated at a sacrificial Zn or Mg anode in an undivided cell, which favor the electrochemical reduction process by decreasing the reduction potential of ArCH(CH3)Cl [14,20]. The electrosyntheses of aryl-2 propionic acids from ArCH(CH3)Cl in the presence of CO2 are efficiently catalyzed by a NiII complex, NiCl2(dppp) (dppp: 1,3-diphenylphosphinopropane Ph2P– (CH2)3–PPh2) associated with the coligand COD (1,5-cyclooctadiene), as reported by Fauvarque et al. [25–29] (Scheme 36.4). Lithium oxalate is oxidized to CO2 at a Ti anode in a one-compartment cell. Zinc powder that is oxidized at the Ti anode can be used as an alternative to lithium oxalate or sacrificial anode [28,29]. The electrosynthesis of fenoprofen has been scaled up to one kilogram (Scheme 36.4) in a specially designed Grignard-type reactor, an undivided cell equipped with a titanium anode and a graphite woven cathode, assembled in a concentric manner to minimize the ohmic drop. The electrolysis is performed at constant current in the nontoxic solvent tetramethylurea (TMU) with 2 mol% of the catalyst NiCl2(dppp) [29]. CO2 is bubbled into the cell at atmospheric pressure and at 0°C to increase its solubility.

ArCH(CH3)Cl + CO2 + 2e 10–50 mmol

NiCl2(dppp) 2 mol% COD 2 mol% TMU, 0°C cathode C, Ti anode

1 atm

CO2H

CO2H iBu CH3

SCHEME 36.4

ArCH(CH3)CO2– + Cl–

O CO2H

CH3 CH3O

CH3

ibuprofen

naproxen

fenoproxen

80%

81%

1 kg, 85%

Nickel-catalyzed electroreductive carboxylation of α-methylbenzyl chlorides.

© 2016 by Taylor & Francis Group, LLC

1398

Organic Electrochemistry

The mechanism of these electroreductive carboxylations is similar to the mechanism established for aryl halides (Scheme 36.3). ArCH(CH3)–NiCl(dppp) generated in the oxidative addition of the electrogenerated Ni0(dppp) to ArCH(CH3)Cl is more easily reduced than ArCH(CH3)Cl. Its reduction by one electron gives ArCH(CH3)–NiI(dppp) which activates CO2. The reaction proceeds via the carboxylation of ArCH(CH3)–NiI(dppp) and not from the anion ArCH(CH3)−. This should allow asymmetric electrosyntheses, in the presence of a nickel catalyst ligated by a chiral ligand. However, up to now the asymmetric electrosynthesis of biologically active chiral aryl-2 propionic acids remains a challenge.

C.

ELECTROREDUCTIVE HOMOCOUPLING OF ARYL/VINYL HALIDES CATALYzED BY NICKEL

1. Electroreductive Homocoupling of Aryl Halides Catalyzed by Nickel The complex NiCl2(dppe), efficient for the electroreductive carboxylation of aryl bromides (see Section II.A), is also a good catalyst for the electrosynthesis of symmetrical biaryls via the electroreductive homocoupling of aryl bromides (Scheme 36.5) [30]. The mechanism of the homocoupling investigated by means of electrochemical techniques is depicted in Scheme 36.6 (paths A–E). After the monoelectronic reduction of ArNiIIBr(dppe) formed in a first oxidative addition (path A), the ensuing complex ArNiI(dppe) undergoes an oxidative addition to ArBr (path C, k2′oa) leading to a NiIII complex, Ar2NiIIIBr(dppe), whose reductive elimination (k3′re) gives the biaryl ArAr and a NiI complex. The latter is reduced back to the initial Ni0 by NiCl2(dppe) 10 mol%

2 ArBr + 2e 10 mmol

SCHEME 36.5

THF + HMPA, rt cathode Hg // Li anode

ArAr + 2Br– 75%

Electroreductive homocoupling of aryl bromides catalyzed by NiCl2(dppe). NiCl2(dppe) –Cl–

1e

(R1)

NiCIL2 –Cl–

1e

(R2)

ArCO2– 2

1e X–

E΄ NiIL

ArCO2

2

E

Homocoupling (A–E) 5 –1 –1 koa 1 = 1.1 × 10 M s –1 oa k΄2 = 960 M s–1 –1 k΄re 3 = 18 s

A

NiIXL2 ArAr

Carboxylation (A,B,C΄–E΄) 5 –1 –1 koa 2 = 1.5 × 10 M s

ArNiIXL2

Reductive elimination k΄re D 3

Ar–NiIIIL2

Ar2NiIIIXL2

O

O

Oxidative addition

koa 2 CO2

1e

B

k΄oa 2 C

ArX

SCHEME 36.6

Oxidative addition

koa 1

1e

D΄

C

THF + HMPA, 20°C L2 = dppe ArX = PhBr

ArX

Ni0L

(R3)

X–

C΄ ArNiIL

2

Mechanism of the electroreductive homocoupling of aryl halides catalyzed by NiCl2(dppe).

© 2016 by Taylor & Francis Group, LLC

Organometallic Compounds as Tools in Organic Electrosynthesis

1399

a one-electron transfer. The evolution of the catalytic reduction plateau current of PhNiIIBr(dppe) (measured at an RDE) with PhBr concentrations allows the determination of the rate constants k2′oa and k3′re (Scheme 36.6). The rate-determining step of the homocoupling is the second oxidative addition (k2′oa) at low PhBr concentrations, whereas the reductive elimination (k3′re) is rate determining at high PhBr concentrations (C > 0.02 M) [30,31]. If the reaction is performed in the presence of CO2, the homocoupling might compete with the carboxylation due to a common intermediate ArNiI(dppe). Scheme 36.6 exhibits the catalytic cycle of the carboxylation (dashed lines, paths A, B, C′–E′) [23,24] in competition with the catalytic cycle of the homocoupling (paths A–E). The two cycles are branched at the level of the key complex ArNiI(dppe), which may react either with CO2 (path C′, k2oa), leading to the carboxylation process, or with ArX in an oxidative addition (path C, k2′oa). The carboxylic acid PhCO2H is formed without any PhPh when the electrolysis is performed from PhBr in the presence of CO2. The carboxylation is thus much more efficient than the homocoupling. Comparison of the values of the two rate constants k2oa and k2′oa confirms that PhNiI(dppe) is indeed ca. 100 times more reactive with CO2 than with PhBr at identical concentrations (Scheme 36.6) [23,30,32]. This is the first reported mechanism of transition metal–catalyzed reactions for which all chemical and electrochemical steps of the catalytic cycles have been characterized, including the determination of rate constants by means of electrochemical techniques. Interestingly, the catalytic cycles of the carboxylation and homocoupling involve a succession of Ni0, NiII, NiI, NiIII, NiI, and Ni0 complexes as a consequence of monoelectronic transfers. The electroreductive homocoupling of aryl halides including the less reactive aryl chlorides is also catalyzed by NiIIBr2(bpy) (bpy = 2,2′-bipyridine) (Scheme 36.7). The electrolyses are performed at room temperature in an undivided cell equipped with a sacrificial Mg anode [33]. The homocoupling is in competition with the reduction, leading to ArH, when the aryl chloride has a substituent in the ortho position [33]. The electrolyses initially performed in DMF or N-methyl-2-pyrrolidone (NMP) have been then performed in protic solvents: alcohols [34], water [35], or ionic liquids (Fe anode [36,37]). The homocoupling has been extended to substituted-2-bromopyridines (Ni foam cathode, Zn or Fe anode) by Navarro et al. [38,39]. The mechanism investigated by Devaud et al. evidences two pathways [40]. Paths A–E in Scheme 36.8 are reminiscent of that established for NiCl2(dppe) (Scheme 36.6). The complex NiBr2(bpy) gives in DMF a mixture of complexes including cationic ones, leading to badly defined close reduction peaks. Nevertheless, the overall bielectronic reduction gives Ni0(bpy), which undergoes oxidative addition to ArX whose rate constant k1oa has been determined (Scheme 36.8). ArNiIIX(bpy) formed in the oxidative addition is reduced by one electron, giving ArNiI(bpy), which undergoes an oxidative addition to ArX to form Ar2NiIIIX(bpy). The rate constant k2oa has been determined (Scheme 36.8). A reductive elimination gives the biaryl ArAr and NiIX(bpy), which is recycled back to the active Ni0(bpy) by a one-electron reduction at the cathode. According to the mechanism, the activation of intermediate ArNiIIX(bpy) by one electron is required and the electrolyses are carried out at the reduction potential of this complex (R2) [40]. An alternative mechanism is also proposed (Scheme 36.8, paths A–B′–C′ or D′) in which ArNiIIX(bpy) gives ArNiIIAr(bpy) and NiIIX2(bpy) after a scrambling of ligands. A reductive

2 ArX + 2e 12 mmol

NiBr2(bpy) 7 mol% 2 bpy 14 mol% NMP or DMF, rt cathode Au, Mg anode

ArAr + 2X– 70–90%

X = I, Br, Cl Ar = Z–C6H4 (Z = CN, COMe, H)

SCHEME 36.7 Electroreductive homocoupling of aryl halides catalyzed by NiBr2(bpy).

© 2016 by Taylor & Francis Group, LLC

1400

Organic Electrochemistry [NiIIBr2(bpy)] –2Br–

X–

2e

(R1) ArX

1e

Ni0(bpy)

koa 1

E NiIX(bpy) ArAr

C΄

Reductive elimination

X–

A

1e D΄

Ar2NiIIIX(bpy)

–1 –1 koa 1 = 69 M s –1s–1 koa = 13 M 2

Oxidative addition

(R1)

1/2 ArAr

D

1/2 NiIIX2(bpy) ArNiIIX(bpy)

1/2 ArNiIIAr(bpy)

Ligand B΄ scrambling

C Oxidative addition

DMF, 20°C bpy = 2,2΄-bipyridine ArX = PhBr

(R2)

1e koa 2

B X–

ArNiI(bpy)

ArX

SCHEME 36.8 Mechanism of the electroreductive homocoupling of aryl halides catalyzed by NiBr2(bpy).

elimination from ArNiIIAr(bpy) gives the biaryl and Ni0(bpy), whereas NiIIX2(bpy) is recycled back to Ni0(bpy) by reduction at the cathode [40]. In this mechanism, the activation of an intermediate organonickel compound by electron transfer is not required and the electrolyses could be carried out at the reduction potential of the NiIIX2(bpy) precursor (R1) instead of that of ArNiIIX(bpy) (R2) in the former mechanism. It is, however, difficult to discriminate between the two mechanisms, since the reduction potentials of ArNiIIX(bpy) and NiIIX2(bpy) may be close to each other, depending on the nature of the aryl group [40]. NiBr2 associated with 2,2′-dipyridylamine (1:1) is much more efficient than NiBr2(bpy) for the electroreductive homocoupling of aryl bromides in ethanol (Fe anode), as reported by Troupel et al. [41]. Interestingly, the electrocarboxylation of aryl halides is not catalyzed by nickel complexes ligated by bipyridine-type ligands, in contrast to nickel complexes ligated by phosphine ligands that catalyze both their carboxylation and their homocoupling, highlighting the key role of the ligand. 2. Electroreductive Homocoupling of Vinyl Halides Catalyzed by Nickel The electroreductive homocoupling of vinyl bromides or chlorides is catalyzed by NiBr2(bpy), leading to 1,3-dienes (Scheme 36.9) [33,42,43]. The electrolyses are carried out in an undivided cell equipped with a sacrificial Al anode in DMF (Scheme 36.9). The mechanism is reported by Labbé et al. [42] (Scheme 36.10). The vinyl halide undergoes an oxidative addition to the electrogenerated Ni0(bpy), leading to vinyl-NiIIX(bpy) (determination of k1oa). The monoelectronic activation of vinyl-NiIIX(bpy) gives vinyl-NiI(bpy) whose oxidative addition to R 2

+ 2e R΄

X

X = I, Br, Cl

NiBr2(bpy) 10 mol% DMF, 20°C cathode Ni, Al anode

R R΄ + Isomers + 2X–

R΄ 50–76%

R

R = H; R΄= Ph, Cy (cyclohexyl) R = Ph, Cy; R΄= H

SCHEME 36.9

Electroreductive homocoupling of vinyl halides catalyzed by NiBr2(bpy).

© 2016 by Taylor & Francis Group, LLC

1401

Organometallic Compounds as Tools in Organic Electrosynthesis [NiIIBr2(bpy)] –2Br–

X– 1e

2e

vinyl-X

Ni0(bpy)

NiIX(bpy)

E

DMF, 20°C Me vinyl-X =

(R1)

Br koa 1

3 –1 –1 koa 1 = 2.5 × 10 M s

Oxidative addition

–1 –1 koa 2 = 55 M s

A

1,3-diene

–1 kre 3 = 2.5 s

kre 3

Reductive elimination

Me

D vinylNiIIX(bpy)

vinyl2NiIIIX(bpy)

koa 2

Oxidative addition

1e B

(R2)

C vinylNiI(bpy)

vinyl-X

X–

SCHEME 36.10 Mechanism of the electroreductive homocoupling of vinyl halides catalyzed by NiBr2(bpy).

the vinyl halide generates (vinyl)2NiIIIBr(bpy) (determination of k2oa). A reductive elimination (determination of k3re) gives the 1,3-diene together with NiI(bpy), which is recycled to Ni0(bpy) by a oneelectron reduction at the cathode. An alternative route involving scrambling ligands in vinyl-NiII complexes (similar to paths A–D′ for aryl halides in Scheme 36.8) is found to be considerably slower than paths A–E in Scheme 36.10 [42].

D.

ELECTROREDUCTIVE POLYMERIzATION OF ORGANIC DIHALIDES CATALYzED BY NICKEL

NiCl2(dppe) catalyzes the electrosynthesis of poly(1,4-phenylene) from 1,4-dibromobenzene (Scheme 36.11) [44,45] or from 4,4′-dibromobiphenyl [46,47]. Similarly, NiBr2(bpy) catalyzes the electroreductive polymerization of 3,6-dibromo-N-substituted carbazole [48], of 3-substituted-2,5-dihalothiophenes (Scheme 36.12) [49]. Ni(bpy)32+ catalyzes the electroreductive polymerization of 2,5-dibromofurane [50] (Scheme 36.13). NiCl2(dppe) 10 mol%

Br

1/n THF + HMPA, rt cathode Hg or C // Li anode

20 mmol

SCHEME 36.11

(

+ 2e

Br

(n

+ 2Br–

n = 4, 5, 6

Nickel-catalyzed electroreductive polymerization of 1,4-dibromobenzene. R

R + 2e

X

S 5 mmol

X

NiBr2(bpy) 10 mol% DMF, rt cathode Ni, Al anode

1/n

(

S

)

+ 2X–

n

25–45%

X = Br, R = H, Me, hexyl, CH2CO2Et X = Cl, R = COMe,

SCHEME 36.12

Nickel-catalyzed electroreductive polymerization of 2,5-dihalothiophene.

© 2016 by Taylor & Francis Group, LLC

1402

Organic Electrochemistry Ni(bpy)32+ 10 mol%

+ 2e Br

Br

O

+ 2Br–

1/n

Acetonitrile, rt cathode Pt, Pt anode

n

O

SCHEME 36.13 Nickel-catalyzed electroreductive polymerization of 2,5-dibromofurane. I

NiI2(dppe) 5 mol%

I + 2e

1/n

n

DMF, rt cathode Pt // Pt anode

SCHEME 36.14

Nickel-catalyzed electroreductive polymerization of diodocacetylene.

Cl

+ 2e

NiCl2(dppe) 10 mol% or NiCl2(dppp) 10 mol%

1/n

DMF–THF, rt cathode Hg // Li anode

Cl 4 mmol

SCHEME 36.15

+ 2I–

+ 2 Cl– n PPX (81–98%)

Nickel-catalyzed electroreductive polymerization of 1,4-Bis(chloromethyl)benzene.

NiI2(dppe) catalyzes the electroreductive polymerization of diiodoacetylene to polyynes (Scheme 36.14) [51]. The electrosynthesis of poly-p-xylene (PPX) is performed via the electrochemical reduction of 4-BrCH2–C6H4 –CH2Br (1,4-bis(bromomethyl)benzene), which is, however, a lachrymatory reagent [52]. The electrosynthesis cannot be performed from the cheaper and less lachrymatory 1,4-bis(chloromethyl)benzene whose reduction potential is much more negative than that of the dibromide derivative. The electroreductive polymerization of 1,4-bis(chloromethyl)benzene is efficiently catalyzed by NiCl2L2 (L2 = dppe, dppp) (Scheme 36.15) via the formation of 4-ClCH2– C6H4 –CH2–NiIIClL2, which are both more easily reduced than 4-ClCH2–C6H4 –CH2Cl [53]. The electrolyses are performed at the controlled reduction potential of 4-ClCH2–C6H4 –CH2– NiIIClL2 determined by cyclic voltammetry. The mechanism of the formation of PPX is similar to that reported for the Ni-catalyzed electroreductive homocoupling of aryl halides (Scheme 36.6, paths A–E), involving an alternation of chemical steps and monoelectronic transfers.

E.

ELECTROREDUCTIVE HETEROCOUPLING OF TWO ORGANIC HALIDES CATALYzED BY NICKEL

The transition metal–catalyzed electroreductive heterocoupling of two different organic halides, RX and R′X′, in order to generate RR′ is very challenging because in competition with the respective homocoupling of both reagents (vide supra). The heterocoupling of aryl halides and α-bromo- or chloroesters catalyzed by NiBr2(bpy) gives ArCH(R)CO2Et (Scheme 36.16), as reported by Sibille et al. [54] and Durandetti et al. [55,56]. The electrosyntheses are performed in an undivided cell equipped with a sacrificial Zn or Al anode. They must be performed upon continuous slow addition of the α-haloester via a syringe pump [54–56] to avoid its competitive oxidative addition to the electrogenerated Ni0(bpy), the latter being much faster than the oxidative addition of Ni0(bpy) to the aryl halide [42]. NiBr2(bpy) 5–10 mol% X΄CH(R)CO2Et + ArX + 2e from 0.3 to 10 mmol

10 mmol

DMF, rt cathode C, Al anode

ArCH(R)CO2Et + X΄– + X– 51–85%

R = H, CH3 X΄= Br, Cl; X = I, Br

SCHEME 36.16 halides.

Nickel-catalyzed electroreductive heterocoupling of α-bromo or α-chloroesters and aryl

© 2016 by Taylor & Francis Group, LLC

1403

Organometallic Compounds as Tools in Organic Electrosynthesis NiBr2(bpy) 5–10 mol%

ClCH(R)COR΄ + ArX + 2e from 0.3 to 10 mmol 10 mmol

DMF, rt (X = I); 70°C (X = Br) cathode Ni, Zn or Al anode

ArCH(R)COR΄ + Cl– + X– 40–80%

R = H, CH3; R΄= Ph, CH3 X = I, Br

Nickel-catalyzed electroreductive heterocoupling of α-chloroketones and aryl halides.

SCHEME 36.17

O

O Cl

R* + ArBr + 2e

Me from 0.3 to 10 mmol 10 mmol

NiBr2(bpy) 10 mol% Ar DMF, rt or 60°C cathode Ni, Al anode

R* =

N (R) Me

Me H2O O

O Me

R* + Cl– + Br–

de 95% S major

Ar

OH

F

N (S) Ph

Ar =

de 93% S major

Me

MeO

SCHEME 36.18 Asymmetric induction in the nickel-catalyzed electroreductive heterocoupling of an α-chloropropionic imide and aryl bromides.

The same group has used the same strategy (slow introduction of the more reactive organic halide) to develop electroreductive heterocoupling of aryl halides with α-chloroketones (Scheme  36.17), benzyl, allyl, or vinyl halides [55,56]. 2- and 3-Bromothiophenes undergo electroreductive heterocoupling with α-chloroesters and benzyl or vinyl halides [57]. The Ni-catalyzed electroreductive heterocoupling of aryl bromides with an α-chloropropionic imide bearing a chiral auxiliary group gives, after hydrolysis, 2-aryl propionic acids, (S)-naproxen, and (S)-flurbiprofen that are anti-inflammatory agents, with good diastereomeric and enantiomeric excesses (Scheme 36.18) [58]. NiBr2(bpy) catalyzes the electrosynthesis of ketones via the electroreductive heterocoupling of aryl halides and benzyl chlorides in the presence of Fe(CO)5 as a CO source (Scheme 36.19) [59]. NiBr2(bpy) catalyzes the electroreductive heterocoupling of aryl halides with 2-halopyridines [60,61] (Equation 1 in Scheme 36.20). It also catalyzes the same reaction with 2-chloropyrimidine [62] and 2-chloropyrazine [62], provided the presence of a catalytic amount of FeBr2 generated at the beginning of the electrolyses by reduction of 30 mol% of 1,2-dibromoethane in the presence of Fe2+ released by a sacrificial Fe anode ((Equations 2 and 3 in Scheme 36.20). In those cases, once the reaction is over, the cell can be charged again with the same amount of ArX and 2-chloropyrimidine or 2-chloropyrazine. The second run gives the same yield as the first one [62]. NiBr2(bpy) catalyzes the electroreductive heterocoupling of aryl halides with 2-chloropyridazines (Scheme 36.21) [63,64]. A mechanistic approach by Barhdadi et al. establishes that the oxidative addition of Ni0(bpy) takes place first with the more reactive 2-chloropyridazine [65]. NiBr2(bpy) 30 mol% ArX + RCH2CI + Fe(CO)5 + 2e 20 mmol 10 mmol 3 mmol

DMF,rt cathode Ni, Fe/Cr/Ni anode

ArCOCH2R + X– + CI– 32–88%

Ar= p–Z–C6H4 (Z = CN, COMe, H, OMe, OH, NMe2) X = I, Br R = p–Z–C6H4 (Z = H, CI); alkyl

SCHEME 36.19 Nickel-catalyzed electrosynthesis of ketones from aryl halides and benzyl chlorides with Fe(CO)5 as the CO source.

© 2016 by Taylor & Francis Group, LLC

1404

Organic Electrochemistry N +

ArBr

NiBr2(bpy) 13 mol%

+ 2e

X

N Ar

DMF/pyridine, 60°C cathode Ni, Fe or Zn anode

(1)

+ Br– + CI–

(2)

31–79%

7.5 mmol

7.5 mmol

+ Br– + X–

Ar = p–Z–C6H4 (Z = CN, CO2Et, COMe, CF3, SO2Me, CI) X = Br, CI 1) FeBr2 30 mol% N ArBr

+ CI

+ 2e N

2 × 3.75 mmol

N

2) NiBr2(bpy) 26 mol% DMF/pyridine, 60°C cathode Ni, Fe anode

Ar N 31–79%

2 × 3.75 mmol

Ar = Z–C6H4 (Z = CN, CO2Et, COMe, CF3, CI, MeO) 1) FeBr2 30 mol%

N ArBr

+ 2e

+ CI

N + Br– + CI–

Ar

DMF/pyridine, 60°C cathode Ni, Fe anode

N 2 × 3.75 mmol

2) NiBr2(bpy) 26 mol%

(3)

N

2 × 3.75 mmol

61%

Ar = p–EtOCO–C6H4

SCHEME 36.20 Nickel-catalyzed electroreductive heterocoupling of aryl bromides and heteroaryl halides.

ArX

+ CI

R + 2e N

5 mmol

N

NiBr2 (bpy) 10 mol% DMF, 60°C cathode Ni, Fe anode

5 mmol

R + X– + CI–

Ar N

N

18–83%

A r = Z–C6H4 (Z = CN, CO2Me, COMe, CF3, F, H, OMe) X = I, Br, CI R = OMe, Me

SCHEME 36.21

Nickel-catalyzed electroreductive heterocoupling of aryl halides and 2-chloropyridazines.

The electroreductive heterocoupling of mono- and dihalopyridines is catalyzed by NiBr2 or by NiBr2(bpy) in the presence of a large excess of NaI, leading to unsymmetrical 2,2′-bipyridines or to terpyridine (tpy). However, similar products are obtained without any nickel catalyst when using a Zn anode instead of a Fe anode [66].

F. ELECTROREDUCTIVE HETEROCOUPLING OF ARYL HALIDES WITH VINYLIC/ ALLYLIC ACETATES OR ARYL HALIDES CATALYzED BY COBALT The electroreductive heterocoupling of aryl halides (ArBr, ArCl) and vinylic acetates is catalyzed by CoBr2 associated with the bpy ligand used in large excess, as reported by Gosmini et al. (Scheme 36.22) [67]. The electrolyses are performed in an undivided cell equipped with a sacrificial Fe anode. The mechanism has been investigated by Buriez et al. (Scheme 36.23, the bpy ligand is voluntarily omitted in all complexes) [68,69]. In paths A–E, the oxidative addition of ArBr to the electrogenerated CoI gives ArCoIIIBr2, which is reduced in two successive one-electron steps (R1 and R2) to ArCoI. ArCoI undergoes an oxidative addition to the vinylic acetate to form ArCoIII(OAc)(η1-vinyl). A reductive elimination generates the coupling product and CoIX. In the case of the less reactive ArCl, an alternative path, A′–B′, takes place (Scheme 36.23). A more reactive electrogenerated Co0

© 2016 by Taylor & Francis Group, LLC

1405

Organometallic Compounds as Tools in Organic Electrosynthesis R ArX +

OAc + 2e

MeCN/pyridine, rt cathode Fe/Cr/Ni, Fe anode

10 mmol 25 mmol X = Br, CI;

R

CoBr2 10 mol% bpy 100 mol%

Ar + X– + AcO– R = H, 40–62% R = Me, 20–92%

R = H, Me

SCHEME 36.22 Cobalt-catalyzed electroreductive heterocoupling of aryl halides and vinyl acetates. CoIIX2 –X–

Ar

1e

(R1)

–

+ AcO Reductive elimination

X–

ArX (X = Br)

CoIX

E Ar

Oxidative addition

A΄

AcO CoIII

(R2) 1e A

Oxidative addition

X–

OAc Co0

D

ArCoIIIX2

OAc

OAc

Co0

ArCoI ArX (X = CI)

C X–

1e

B

(R1)

Oxidative addition ArCoIIX

1e

SCHEME 36.23 a vinyl acetate.

B΄

X–

(R2)

Mechanism of the cobalt-catalyzed electroreductive heterocoupling of aryl bromides and

species, stabilized by complexation to the C=C bond of the vinylic acetate, undergoes an oxidative addition to ArCl to give ArCoIIX. The complexes CoIX and ArCoIIX are reduced at the same potential (R2), which is more negative than the common reduction potential of CoIIX2 and ArCoIIIX2 (R1). The electrolyses are efficient, provided they are carried out at the reduction potential of ArCoIIX (R2) and in the presence of cations Fe2+ released in the oxidation of a sacrificial Fe anode. Fe2+ cations play a key role by trapping the acetate ions responsible for the catalyst poisoning [68,69]. The electroreductive heterocoupling of aryl halides (ArBr, ArCl) and allyl acetate or carbonate is catalyzed by CoBr2, as reported by Gosmini et al. (Scheme 36.24) [70]. The pyridine used as a cosolvent also serves as a ligand for Co complexes. The electrolyses are performed in an undivided ArX + 7.5 mmol 20 mmol

OY

+ 2e

CoBr2 13 mol% MeCN / pyridine, 50°C cathode Fe/Cr/Ni, Fe anode

X = Br, Cl Y = COMe, CO2Me Ar = p-Z–C6H4 (Z = COMe, CO2Et, CF3, CN, F, H, OMe )

Ar

+ X– + YO–

Y = COMe, 30–74% Y = CO2Me, 26–50%

SCHEME 36.24 Cobalt-catalyzed electroreductive heterocoupling of aryl halides and allylic acetate or carbonate.

© 2016 by Taylor & Francis Group, LLC

1406

Organic Electrochemistry

I + 2e

Br + Z΄

Z 10 mmol

CoBr2 30 mol%

+ Br– + I–

MeCN / pyridine, 20°C cathode Fe/Cr/Ni, Fe anode Z

more reactive 20–30 mmol

Z΄ 57–90%

o or p-Z = CO2Et, CO2Me, CN o,m or p-Z΄ = OMe, CF3, H

SCHEME 36.25 Cobalt-catalyzed electroreductive heterocoupling of aryl bromides and aryl iodides.

cell equipped with a sacrificial Fe anode. The mechanisms proposed by Buriez et al. for allyl acetate [71] or allyl ether [72] are very close to that established for vinyl acetate (Scheme 36.23). In a competitive process, CoBr2/pyridine catalyzes the electroreductive homocoupling of allylic acetates to generate 1,5-hexadienes [73]. CoBr2 associated with pyridine catalyzes the electroreductive heterocoupling of two different aryl halides (ArX and Ar′X′) to deliver unsymmetrical biaryls ArAr′ (Scheme 36.25) [74]. The symmetrical ArAr and Ar′Ar′ may be generated as by-products. The reactions are regiospecific in all cases.

III.

TWO-ELECTRON ACTIVATION OF ORgANOMETALLIC COMPLEXES IN TRANSITION METAL–CATALyZED ELECTROSyNTHESES

A.

ELECTROREDUCTIVE CARBOXYLATION OF ARYL HALIDES/TRIFLATES CATALYzED BY PALLADIUM

1. Electroreductive Carboxylation of Aryl Halides Catalyzed by Palladium The palladium-catalyzed electroreductive carboxylation of aryl iodides and bromides has been reported by Torii et al. [75]. Pd0(PPh3)4 is used as a catalyst or PdCl2(PPh3)2 which is reduced to Pd0 at the very beginning of the electrolyses. The reactions are regiospecific (Scheme 36.26). A mechanistic approach has established that the isolated trans-PhPdBr(PPh3)2, formed in the oxidative addition of Pd0(PPh3)2 (generated in situ from Pd0(PPh3)4) to PhBr, is reduced at less negative potential than PhBr. Its reduction in the presence of CO2 gives the desired carboxylic acid PhCO2H [75]. The mechanism of the carboxylation step has been more deeply investigated by Amatore et al. by using electrochemical techniques [76]. The complexes ArPdX(Cl)(PPh3)2− [77] or trans-ArPdX(PPh3)2 (X = I, Br), formed in the oxidative addition of ArX to electrogenerated Pd0(PPh3)2Cl− or Pd0(PPh3)2, respectively, are reduced at less negative potentials than the parent ArX (Scheme 36.27) [76]. Their reduction in an overall two-electron process generates the anionic ArPd0(PPh3)2− in equilibrium with Pd0(PPh3)2 and the anion Ar− (Scheme 36.26, paths A–B–C or A′–B′–C) [76]. The carboxylation proceeds classically from Ar−, that is, outside the coordination sphere of the palladium, from the same Ar−, which would have been generated by the electrochemical reduction of ArX (whenever possible). Importantly, Ar− is generated by the reduction of ArPdX(Cl)(PPh3)2− or ArPdX(PPh3)2, that is, at a less negative potential than the reduction potential of ArX (X = I, Br) [76]. Consequently, aryl halides (such as 4-MeO–C6H4–Br), which are reduced at very negative potentials (beyond that of CO2), can be PdCl2(PPh3)2 7 mol% or Pd0(PPh3)4 7 mol%

ArX + CO2 + 2e 0.3 mmol 1 atm

ArCO2– + X– DMF, rt 47–92% cathode Pt or Pb // Pt anode

X = I, Br Ar = p-Z–C6H4 (Z = Cl, H, Me, OMe, tBu)

SCHEME 36.26 Palladium-catalyzed electroreductive carboxylation of aryl halides.

© 2016 by Taylor & Francis Group, LLC

1407

Organometallic Compounds as Tools in Organic Electrosynthesis Pd0(PPh3)4

PdCl2(PPh3)2 – Cl–

THF, 20°C L = PPh3 ArX = PhI

2e

–L

(R1) ArX

Pd0L2Cl– or Pd0L2X–

Cl–

–1 –1 koa 1 = 530 M s

or X–

D

Pd0L3 +L

Oxidative addition

koa 1

–L

Pd0L2 Pd0L2

ArX Oxidative L addition

A A΄ L Ar

Pd

X

– C l–

Cl

Ar

Pd

L –

ArCO2

CO2

X

L

Ar– C

B 2e (R2)

Ar

Pd0L2–

B΄

Cl–, X–

2e (R3) X–

SCHEME 36.27 Mechanism of the palladium-catalyzed electroreductive carboxylation of aryl halides.

transformed into the corresponding carboxylic acids, via the easier reduction of aryl-PdII complexes, at the origin of the catalysis by palladium complexes [76]. In contrast to the electroreductive carboxylation catalyzed by nickel ligated to dppe that involves only monoelectronic transfers (Scheme 36.3), the catalytic cycle of the electroreductive carboxylation catalyzed by palladium ligated to PPh3 only involves Pd0 and PdII complexes as a consequence of bielectronic transfers, as evidenced by the determination of the number of electron(s) involved in each step [4]. Therefore, two different kinds of activation of ArMXL2 (M = Ni, Pd) complexes are effective: a one-electron transfer for the Ni complex ligated by the bisphosphine dppe (Scheme 36.3) [24] vs. an overall two-electron transfer for the Pd complex ligated by the monophosphine PPh3 (Scheme 36.27) [76]. This affects the mechanism of the carboxylation step: within the Ni coordination sphere in the former case (Scheme 36.3) vs. outside the coordination sphere in the Pd case (Scheme 36.27), highlighting the crucial role of the metal and ligand. 2. Electroreductive Carboxylation of Aryl Triflates Catalyzed by Palladium Aryl triflates (Ar–OTf: ArOSO2CF3) are easily synthesized from widely available phenols. They are often used as an alternative to aryl halides [78]. The OTf group could be considered as a better leaving group than halides. However, even if most aryl triflates are reduced at a cathode in DMF in one bielectronic step, an electrolysis at their reduction potential in the presence of CO2 generates the phenols as the major product and not the desired aromatic carboxylic acids (Scheme 36.28), as reported by Jutand and Négri [79]. The electrochemical reduction of an aryl triflate first generates a radical anion ArOTf •− whose major evolution is not the cleavage of the Ar–O bond, which would generate (after a second electron transfer taking place at the potential of the fist one) the aryl anion Ar− prone to react with CO2, but the cleavage of the O–S bond leading to the undesired phenoxide after a second electron transfer (Scheme 36.28) [79]. This problem has been bypassed by Jutand et al., by using a palladium(0) catalyst that selectively cleaves the Ar–O bond and not the O–S bond (Scheme 36.29) [79,80]. Indeed, whereas the electrochemical reduction of 1-naphthyl triflate in the presence of CO2 leads, after hydrolysis, to only 13% of the 1-naphthyl carboxylic acid (Scheme 36.28),

© 2016 by Taylor & Francis Group, LLC

1408

Organic Electrochemistry DMF, rt ArOTf + CO2 + 2e ArCO2– + ArO– + TfO– + Tf– cathode C // Mg anode 1 mmol 1 atm 13% 80% Ar = 1–naphthyl Mechanism –

Ar–OTf

Ar–OTf + 1e

– T fO– Ar–O–SO2CF3–

Ar

– T f–

Ar–O

1e 1e

Ar–

CO2

ArCO2–

Ar–O–

Major path

SCHEME 36.28

Electrochemical reduction versus electrocarboxylation of aryl triflates. PdCl2(PPh3)2 10 mol% ArOTf + CO2 + 2e 1 mmol 1 atm

DMF, 90°C cathode C // Mg anode

ArCO2– + TfO– 22–96%

Ar = p-Z–C6H4 (Z = CN, CF3, CH3CO, CO2Et, Br, Cl, F, H, CH3, tBu); 1-naphthyl; 2-naphthyl; 2-pyridyl

SCHEME 36.29

Palladium-catalyzed electroreductive carboxylation of aryl triflates.

addition of 10 mol% of PdCl 2(PPh3)2 results in the formation of the carboxylic acid in 83% yield (Scheme 36.29). The electrolysis is performed at a less negative potential than that of 1-naphthyl triflate, under CO2 at atmospheric pressure. The reaction has been extended to various aryl triflates. It is regiospecific and exhibits good compatibility with substituents Z on the aryl group (Scheme 36.29) [79,80]. As an application, the 1,3,5-tris(4-carboxyphenyl)benzene (3-CO2H) has been synthesized via the Pd-catalyzed electroreductive carboxylation of 3-OTf (Scheme 36.30) [79]. In the absence of CO2, an electrolysis performed at the reduction potential of 3-OTf without any catalyst delivers 1,3,5-tris(4-OH-phenyl)benzene (3-OH) (Scheme 36.30). Both molecules are precursors of materials OTf

CO2H

+ CO2 + 6e

1) PdCl2(PPh3)2, 30 mol% DMF, 90°C cathode C // Mg anode

1 atm TfO

2) H2O

OTf

3-OTf 1 mmol

+ 3 TfO–

HO2C

3-CO2H 75%

OH 1) 6e DMF, rt + 3 Tf –

2) H2O

HO

SCHEME 36.30

3-OH 67%

OH

Palladium-catalyzed electroreductive carboxylation of aryl tris-triflates.

© 2016 by Taylor & Francis Group, LLC

CO2H

1409

Organometallic Compounds as Tools in Organic Electrosynthesis PdCl2(PPh3)2 – Cl–

2e

DMF, 20°C L = PPh3 Ar = 1-Naphthyl –1 –1 koa 1 = 5.5 M s

(R1) ArOTf

Pd0L2Cl– Oxidative addition

k1oa

Cl–

TfO–

A L Pd0L2

Ar

Pd

Cl

L

ArCO–2

CO2

B

C

2e

Ar– Ar

SCHEME 36.31

Pd0L2–

(R2)

Cl–

Mechanism of the palladium-catalyzed electroreductive carboxylation of aryl triflates. 

exhibiting nonlinear optics properties [81]. The direct catalyst-free electrochemical reduction is a procedure for the deprotection of aryl triflates to phenols [79]. The mechanism of the Pd-catalyzed electroreductive carboxylation of aryl triflates is close to the mechanism of the Pd-catalyzed electroreductive carboxylation of aryl halides except that it involves neutral aryl-PdII complexes (Scheme 36.31) [79]. The anionic Pd0(PPh3)2Cl− is generated at the very beginning of the electrolysis by the bielectronic reduction of the precursor PdCl2(PPh3)2 (R1) [4,77]. Its oxidative addition to ArOTf generates a neutral complex ArPdIICl(PPh3)2, due to the low affinity of the triflate anion TfO− for PdII centers [82]. Determination of the rate constants of oxidative additions reveals the reactivity order: PhI ≫ PhOTf > PhBr [82]. ArPdCl(PPh3)2 is activated by an overall two-electron transfer (R2), which takes place at a less negative potential than the reduction potential of ArOTf (e.g., Ep = −2.2 V for PhPdCl(PPh3)2 and −2.63 V vs. SCE for PhOTf in DMF) [79]. The reduction peak potentials of ArPdCl(PPh3)2 complexes have been determined by cyclic voltammetry to fix the potential of the electrolyses. The  reduction of ArPdCl(PPh3)2 generates the anionic ArPd0(PPh3)2− in equilibrium with the anion Ar− and Pd0(PPh3)2 [76]. The catalytic cycle is closed by coordination of a chloride ion (Scheme 36.31). The carboxylation proceeds from the anion Ar−, outside the coordination shell of the palladium. The electrocarboxylation of aryl triflates is therefore more challenging than that of aryl halides since the electrochemical reduction of aryl halides (whenever possible) generates the desired anion Ar− (Scheme 36.1). The chemical (A) and electrochemical (B) activations in Scheme 36.31 are both crucial. Indeed, the palladium (via the activation of the Ar–O bond by oxidative addition) catalyzes the reduction of ArOTf to the desired Ar− and avoids the undesired formation of ArO − that would be obtained as the major process in the direct electrochemical reduction of ArOTf (Scheme 36.28). Interestingly, the chloride ions delivered during the reduction of the precursor PdCl2(PPh3)2 exert a beneficial effect by providing neutral ArPdCl(PPh3)2 instead of cationic ArPd(DMF)(PPh3)2+ [82]. Indeed, the electrochemical reduction of cationic ArPd(DMF)(PPh3)2+ is a monoelectronic process that gives ArPdI(PPh3)2 and then the biaryl ArAr [83]. As aryl triflates are synthesized from phenols, their palladium-catalyzed electroreductive carboxylation can be compared to the carboxylation of phenoxides (the known Kolbe–Schmitt reaction [84]),

© 2016 by Taylor & Francis Group, LLC

1410

Organic Electrochemistry

O–

OH

CO2

OH Kolbe–Schmitt CO2–

OH

Tf2O

OTf

CO2, 2e [Pd cat]

CO2–

Jutand

SCHEME 36.32 Direct electroreductive carboxylation of phenol versus palladium-catalyzed electroreductive carboxylation of phenol via phenyl triflate.

which leads to carboxylic acids without alteration of the phenol function (Scheme 36.32). The palladium-catalyzed electroreductive carboxylation is the only way to generate aromatic carboxylic acids from phenols with cleavage of the aromatic C–O bond via their easily available triflates ArOTf (Scheme 36.32) [79,80]. Moreover, the electron transfer causes an inversion of reactivity (umpolung). Aryl triflates can now react with an electrophile (CO2), whereas they are usually known to react with nucleophiles in cross-coupling reactions, in the presence of the same PdCl2(PPh3)2 catalyst [78].

B.

ELECTROREDUCTIVE CARBOXYLATION OF VINYL TRIFLATES CATALYzED BY PALLADIUM

Vinyl triflates are easily synthesized by the reaction of triflic anhydride with ketones (via the enols) or triflic acid with alkynes [78]. Jutand and Négri have reported that the direct electrochemical reduction of a vinyl triflate gives, after consumption of two electrons per mole, the corresponding ketone, via the cleavage of the O–S bond (Scheme 36.33) [85]. As reported by Tokuda et al. [86], the direct electrochemical reduction of alkyl-substituted vinyl triflates, performed in the presence of CO2 and Mg2+ cations released by a sacrificial Mg anode, affords the β-keto carboxylic acids (Scheme 36.34), by carboxylation of the enolate generated by the electrochemical cleavage of the O–S bond revealed in Scheme 36.33. According to the authors, this − reduction process might be induced by CO•2 , the reduced form of CO2 that is more easily reduced than alkyl-substituted vinyl triflates in DMF [79,87]. In contrast, the same reactions performed in the presence of PdCl2(PPh3)2 as catalyst are performed at a less negative potential and give the α,β-unsaturated carboxylic acids (Scheme 36.35) [79,85]. The reactions are stereospecific. The electrolyses are performed at room temperature. The reactions are thus faster than the carboxylations of aryl triflates (Scheme 36.29) because oxidative additions of vinyl triflates to OSO2CF3 + 2e (OTf ) 1 mmol

SCHEME 36.33

– Tf – DMF, rt cathode C // Mg anode

O–

H2O 98%

Electrochemical reduction of vinyl triflates. CO2– OTf + CO2 + 2e

DMF, rt cathode Pt, Mg anode

6 mmol

OTf

SCHEME 36.34

1-, 2-naphthyl-OTf

Electroreductive carboxylation of vinyl triflates.

© 2016 by Taylor & Francis Group, LLC

O + Tf– 28–77%

n = 1–3

O

1411

Organometallic Compounds as Tools in Organic Electrosynthesis

PdCl2(PPh3)2 10 mol% OTf + CO2 + 2e 1 mmol

CO2– + TfO–

DMF, rt cathode C // Mg anode

1 atm CO2H

CO2H CO2H

CO2H

CO2H

CO2H

Ph

SCHEME 36.35

60%

32%

85%

80%

70%

86%

Palladium-catalyzed electroreductive carboxylation of vinyl triflates.

Pd0 complexes are faster than those of aryl triflates. The respective rate constants have been determined by means of electrochemical techniques [82,87]. As for aryl triflates, the cleavage of the vinylic C–O bond of vinyl-OSO2CF3 by the electrogenerated Pd0(PPh3)2Cl− catalyst is favored vs. the O–S bond cleavage. The neutral vinyl-PdCl(PPh3)2 complexes formed in the oxidative addition whose reduction potentials have been determined by cyclic voltammetry are more easily reduced than the parent vinyl-OTf [79]. The electrolyses are performed at the reduction potential of vinyl-PdCl(PPh3)2. This avoids the direct reduction of vinyl-OTf to the undesired enolate (Scheme 36.33). The carboxylation might proceed from a vinylic anion generated in the bielectronic reduction of vinyl-PdCl(PPh3)2 by analogy to the carboxylation of aryl triflates (Scheme 36.31) [79]. Consequently, the activation of alkyl-substituted vinylic triflates by the Pd0 catalyst is crucial, since the palladium catalyzes their reduction and moreover favors the desired cleavage of the vinylic C–O bond. Vinyl triflates are synthesized from ketones or alkynes [78], which now can be transformed into α,β-unsaturated carboxylic acids in the presence of CO2, an electron source and a Pd catalyst, via the corresponding vinyl triflates (Scheme 36.36). When vinyl triflates are generated from terminal alkynes, their Pd-catalyzed stereospecific carboxylation selectively affords one α,β-unsaturated carboxylic acid with the carboxylic group at the more substituted position (Schemes 36.35 and 36.36) [79,85]. The Ni-catalyzed electroreductive hydrocarboxylation of terminal alkynes has been reported by Duñach et al. but the reaction is not regioselective and affords a mixture of two α,β-unsaturated carboxylic acids (see Section IV.A.1). Phenyl-substituted vinyl triflates are more easily reduced than alkyl-substituted ones. They are converted to α,β-unsaturated carboxylic acids in the presence of Mg2+, by cleavage of the vinylic C–O bond [88]. Their reduction potentials are nevertheless still very negative. The palladium catalyst makes their reduction proceed at less negative potentials via the formation of the more easily reduced vinyl-PdIICl complexes. The electroreductive carboxylation of vinyl triflates catalyzed by palladium is an alternative route to the electroreductive carboxylation of vinyl halides catalyzed by Mg2+ ions [13,20,89], nickel [90], or palladium [75] complexes.

O

Tf2O or PhNTf2

OTf + CO2 + 2e

Base

HOTf

R

R

CO2– + TfO–

R + CO2 + 2e

OTf

SCHEME 36.36

[Pd] cat

[Pd] cat

+ TfO– CO2–

Synthesis of vinyl triflates and their palladium-catalyzed electroreductive carboxylation.

© 2016 by Taylor & Francis Group, LLC

1412

Organic Electrochemistry R΄ R

OAc OAc

R = Ph R΄ = H

H+ 2e + E

+

ClSiMe3 R΄

R

Pd0(PPh3)4 5 mol%

R

H

45–58%

acetonitrile, rt cathode Pb // Pt anode

R

SiMe3

82–68%

H or SiMe3

0.31 mmol

R΄

R

R = Ph R΄ = Me

R΄

R

Pd0L2

R

2e R΄

R

SCHEME 36.37 electrophiles.

major (2/1)

E

L + L Pd

C.

H: 64% SiMe3: 68%

H or SiMe3

R

CO2

R΄

–Pd0L2

R = Ph R΄ = H

CO2H R

CO2H 40% 36%

Palladium-catalyzed electroreductive cleavage of allylic acetates followed by reaction with

ELECTROREDUCTIVE CLEAVAGE OF ALLYLIC ESTERS CATALYzED BY PALLADIUM OR NICKEL

Allylic acetates usually react with nucleophiles in the presence of a palladium(0) catalyst (Tsuji– Trost reaction) [3]. Their reactivity has been inverted in an electroreductive cleavage catalyzed by Pd0(PPh3)4, which allows their reaction with electrophiles, as reported by Torii et al. (Scheme 36.37) [91]. The key step is the bielectronic reduction of the cationic (η3-allyl)palladium(II) complexes generated in the oxidative addition of the Pd0 catalyst to the allylic acetate. An allylic anion is generated that reacts with electrophiles, leading to the inner alkene rather than the terminal one. This procedure is an electrochemical deprotection of allylic esters (E = H+). CO2 has been used as an electrophile leading to the electrosynthesis of β,γ-unsaturated carboxylic acids (Scheme 36.37) [75]. Allylic acetates or carbonates may similarly react with CO2 in the presence of NiBr2(bpy) in DMF to give β,γ-unsaturated carboxylic acids in a one-compartment cell (C cathode, Mg anode) as reported by Duñach et al. [92].

D.

ELECTROREDUCTIVE HOMOCOUPLING OF ARYL HALIDES/TRIFLATES CATALYzED BY PALLADIUM

1. Electroreductive Homocoupling of Aryl Halides Catalyzed by Palladium The palladium-catalyzed electroreductive homocoupling of aryl halides has been reported by Torii et al. (Scheme 36.38) [93].

2 ArX + 2e 0.3 mmol

PdCl2(PPh3)2 7 mol% or Pd(PPh3)4 7 mol%

ArAr DMF, rt 87–98% cathode Pb // Pt anode

+ 2X–

X = I, Br Ar = p-Z–C6H4 (Z = H, Me, NMe2, OMe, tBu)

SCHEME 36.38

Palladium-catalyzed electroreductive homocoupling of aryl halides.

© 2016 by Taylor & Francis Group, LLC

1413

Organometallic Compounds as Tools in Organic Electrosynthesis Pd0(PPh3)4

PdCl2(PPh3)2 – Cl–

THF, 20°C L = PPh3 ArX = PhI or X–

Reductive elimination D

Pd0L2

L Ar

Pd L

CO2

–L

(R1)

Pd0L3

ArX Pd0L2Cl– or Pd0L2X–

Cl–

–1 –1 koa 1 = 530 M s 4 M–1s–1 koa = 2 × 10 2

ArCO2–

2e

Ar–

E

X

+L

Oxidative addition

koa 1

–L

Pd0L2

ArAr

k1΄oa

A L

– Ar

Ar

Pd L

A΄

– X

L

– Cl– Ar

Cl

ArX Oxidative addition

Pd

X

L

Oxidative addition ArX

B

koa 2

2e (R2)

C Ar

B΄ Pd0L2–

Cl–+ X–

2e (R3) X–

SCHEME 36.39 Mechanism of the palladium-catalyzed electroreductive homocoupling of aryl halides.

The mechanism has been established by Amatore et al. by means of electrochemical techniques (Scheme 36.39, paths A–B–C–D or A′–B′–C–D) [94]. ArX undergoes an oxidative addition to the anionic ArPd0(PPh3)2− leading to the anionic ArPdIIAr(X)(PPh3)2−, which gives the biaryl and a Pd0 complex by reductive elimination. Two oxidative additions are involved (A or A′ and C), separated by a two-electron reduction (B or B′), respectively. The second oxidative addition (k2oa in C) is faster than the first one (k1oa in A) by a factor 40 in the case of PhI (Scheme 36.39) [94]. In the presence of CO2, no biaryl is formed, indicating that the electroreductive carboxylation (dashed lines in Scheme 36.39) is more efficient than the electroreductive homocoupling. The two cycles are branched at the level of ArPd0L2− that can either react with ArX (path C) leading to ArAr or dissociate to Ar− (path E), leading to Ar− and then to ArCO2− (Scheme 36.39). 2. Electroreductive Homocoupling of Aryl Triflates Catalyzed by Palladium PdCl2(PPh3)2 catalyzes the electroreductive homocoupling of aryl triflates to symmetrical biaryls (Scheme 36.40) [80,95]. The reactions are regiospecific. The electrolyses are carried out at the reduction potential of the complexes ArPdCl(PPh3)2 [95]. The mechanism of the homocoupling is displayed in Scheme 36.41 (paths A–B–C–D). The aryl triflate is activated twice: in its oxidative addition to the electrogenerated Pd0(PPh3)2Cl− (A) and in its oxidative addition to the electrogenerated ArPd0(PPh3)2− (C). Interestingly, the Ar–O bond is selectively cleaved in both oxidative additions. The biaryl ArAr was not formed when the electrolyses were performed in the presence of CO2 [79,80]. The two competitive reactions, carboxylation and homocoupling, proceed via a

2 ArOTf + 2e 1–10mmol

PdCl2(PPh3)2 1–10 mol% DMF, 90°C cathode C // Mg anode

ArAr + 50–70%

2TfO–

Ar = p-Z–C6H4 (Z = CN, CF3, Cl, H); 1-naphthyl; 2-naphthyl; 2-pyridyl

SCHEME 36.40

Palladium-catalyzed electroreductive homocoupling of aryl triflates.

© 2016 by Taylor & Francis Group, LLC

1414

Organic Electrochemistry PdCl2(PPh3)2 –Cl–

DMF L = PPh3 (R1)

2e

ArOTf Pd0L2Cl– Cl–

Oxidation addition

Cl– Reductive elimination

TfO–

ArAr A

D

L

L Pd0L2

Ar

Ar

L

Pd

Pd

Cl

L

Ar TfO– Oxidative addition

E ArCO2–

CO2

B

ArOTf

Ar–

2e

(R2)

C Ar

Pd0L–2

Cl–

SCHEME 36.41 Mechanism of the palladium-catalyzed electroreductive homocoupling of aryl triflates.

common intermediate: the anionic ArPd0(PPh3)2− (Scheme 36.41) [94,95]. The carboxylation is much more efficient than the homocoupling.

E.

ELECTROREDUCTIVE HETEROCOUPLING OF TWO ARYL HALIDES CATALYzED BY PALLADIUM

Torii et  al. have explored the challenging Pd-catalyzed electroreductive heterocoupling of two different aryl halides, ArX and Ar′X′ in order to generate dissymmetrical biaryls Ar–Ar′ (Scheme 36.42) [93]. [Pdıı] or [Pd0]

ArAr΄

ArX + Ar΄X΄ + 2e

Catalytic reaction (A) l + 2e

l + tBu

N 0.3 mmol

0.3 mmol

PdCl2(PPh3)2 7 mol% PPh3 14 mol%

tBu

N

DMF, rt cathode Pb//Pt anode

45%

+ tBu

tBu

N

N

21%

tBu

Stoichiometric reaction (B) L = PPh3 –2L N l+ DMF, rt 0.3 mmol 0.3 mmol Pd0L4

SCHEME 36.42

24%

2e N

l 0.3 mmol

N DMF, rt cathode Pb // Pt anode

tBu

PdIL2

77%

Palladium-catalyzed electroreductive heterocoupling of two aryl halides.

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Organometallic Compounds as Tools in Organic Electrosynthesis

1415

An electrolysis of a solution of ArI (4-tBu–C6H4 –I) and Ar′I (4-Me2N–C6H4 –I) in the presence of a palladium catalyst leads to a mixture of ArAr′, ArAr, and Ar′Ar′ (Scheme 36.42 (A)) [94]. In contrast, a reaction performed step by step, using a stoichiometric amount of Pd0(PPh3)4, leads to the major formation of the desired dissymmetrical biaryl ArAr′ (Scheme 36.42 (B)) [93]. ArPdXL2 formed in the oxidative addition of Pd0(PPh3)2 to ArX must be reduced to be able to react with the second aryl halide Ar′X′, in agreement with the mechanism proposed by Amatore et  al. for the Pd-catalyzed homocoupling of aryl halides: two oxidative additions of aryl halides, separated by a two-electron reduction (Scheme 36.39, A′–B′–C) [94]. The mechanism of the heterocoupling of two different aryl halides has been further investigated. The first oxidative addition of Pd0(PPh3)2 to aryl halides follows the decreasing reactivity orders (electron-donating group [EDG]; electronwithdrawing group [EWG]) [96,97]: ArI > ArBr ≫ ArCl 4-EWG–C6H4 –X > 4-EDG–C6H4 –X If at similar concentrations ArX is more reactive than Ar′X′ in the first oxidative addition, the electrogenerated anionic ArPd0L2− will be first formed and could react with either ArX or Ar′X′ in the second oxidative addition. Such a reaction is about 40 times faster than the first oxidative addition (for the same ArX) but is less sensitive to electronic factors (nevertheless with the same reactivity order) [94]. Therefore, the more reactive aryl halide ArX in oxidative additions will favor the formation of ArAr first. ArAr′ will be formed when the concentration of ArX becomes much lower than that of Ar′X′ to allow competition with ArX in the second oxidative addition. From this mechanism, one predicts that a Pd-catalyzed electroreductive heterocoupling performed on a stoichiometric mixture of ArX and ArX′ will lead to a mixture of ArAr, ArAr′, and Ar′Ar′, which has been observed experimentally (Scheme 36.42 (A)) [94]. Such problem may be partly solved if the electrolyses are performed in a two-step batch procedure, each batch involving a stoichiometric amount of palladium as in Scheme 36.42 (B). An alternative is the slow introduction of the most reactive aryl halide during the electrolyses via a syringe pump, so that to compensate its intrinsic higher reactivity in oxidative additions by a lower concentration.

F.

ELECTROSYNTHESIS OF KETONES VIA THE ELECTROREDUCTIVE HETEROCOUPLING OF ORGANIC HALIDES CATALYzED BY NICKEL

The electroreductive heterocoupling of benzyl bromides/allylic chlorides with acyl chlorides gives ketones in the presence of a NiII catalyst: Ni(bpy)3(BF4)2 (Scheme 36.43) [98]. The electrolyses are performed at the reduction potential of the NiII precursor, that is, at a less negative potential than the reduction potential of the two organic halides. The mechanism of the heterocoupling between PhCH2COCl and PhCH2Br has been investigated by Amatore et al. (Scheme 36.44) [99]. The bielectronic one-step reduction of the cationic precursor NiII(bpy)32+ (R1) at the beginning of the electrolysis gives the Ni0(bpy)2 complex [4]. After dissociation of one ligand, Ni0(bpy) RX + R΄COCl + 2e

Niıı (bpy)3(BF4)25 mol % acetonitrile, rt cathode Au or C, Zn or Mg anode

R΄COR + Cl– + X– 35–96%

R = PhCH2, allyl X = B r, C l R΄ = PhCH2, Et, Ph, iPr

SCHEME 36.43 Nickel-catalyzed electrosynthesis of ketones via electroreductive heterocoupling of two organic halides.

© 2016 by Taylor & Francis Group, LLC

1416

Organic Electrochemistry Niıı(bpy)32+ –bpy

2e

acetonitrile, 20 °C bpy= 2,2΄–bipyridine

(R1)

koa = 8 × 105 M–1s–1

Ni0(bpy)2 +bpy

–bpy

PhCH2Br Ni0(bpy)

PhCH2COCH2Ph Reductive elimination

koa

PhCH2Niıı– COCH2Ph(bpy)

Oxidative addition

PhCH2NiııBr(bpy)

2e

(R1)

Cl–

PhCH2COCl

Br–

PhCH2Ni0(bpy)–

SCHEME 36.44 Mechanism of the nickel-catalyzed electroreductive heterocoupling of PhCH2Br and PhCH2COCl.

undergoes oxidative addition with PhCH2Br, leading to PhCH2NiIIBr(bpy). The rate constant of the oxidative addition (koa) to PhCH2Br has been determined by fast cyclic voltammetry and found to be 10 times higher than that of PhCH2COCl. The overall two-electron reduction of PhCH2NiIIBr(bpy) (at the same or at less negative potential than the reduction potential of NiII(bpy)32+ at R1) generates an anionic complex PhCH2Ni0(bpy)−, which reacts with PhCH2COCl by nucleophilic substitution to give a ketone after reductive elimination from PhCH2NiII–COCH2Ph(bpy) [99] (Scheme 36.44). The reductive heterocoupling requires a double activation: chemical activation of the benzylic bromide by a Ni0 complex followed by activation of the resulting PhCH2NiBr(bpy) by a two-electron transfer, in contrast to the one-electron transfer established for ArNiBr(bpy) (Scheme 36.8).

G. ELECTROSYNTHESES OF CYCLOPROPANES CATALYzED BY NICKEL OR COPPER Substituted cyclopropanes are formed in the NiBr2-catalyzed electroreductive coupling of gemdibromo compounds (CH3CHBr2 or CH2Br2) and electron-deficient alkenes, as reported by Léonel et  al. (Scheme 36.45) [100]. Pyridine used as a cosolvent also serves as ligands for the nickel centers. A nickel catalyst is required when the alkene is less easily reduced than the gem-dihalide compound (RCHX2). When the electron-deficient alkene is more easily reduced than the gem-dihalide compound, its reduction in the presence of RCHX2 leads to substituted cyclopropanes with a stereochemistry opposite to that obtained in the Ni-catalyzed reactions [101]. This suggests a change CO2Me + 2e

CH2Br2 + 10 mmol

CO2Me 25 mmol

NiBr220 mol % CH3CN/pyridine, 60°C cathode Ni, Mg anode

+ CO2Me MeO2C

MeO2C

(Z)

65% (cis/trans = 80/20)

(E)

56% (trans)

SCHEME 36.45 Nickel-catalyzed electrosyntheses of cyclopropanes.

© 2016 by Taylor & Francis Group, LLC

+ 2Br– CO2Me

1417

Organometallic Compounds as Tools in Organic Electrosynthesis R

R

EWG

RR΄2CCl2 + 20 mmol

+ 2e

10 mmol

R΄

CuBr 10 mol% DMF/pyridine, –10°C cathode Ni, Fe anode

+ 2 Cl– R

EWG 25–76%

R΄ = R΄ = Ph; R = H, R΄ = Ph; R = Cl, R΄ = Ph or CO2Me R= H, Ph, Me EWG = CO2Me, COMe

SCHEME 36.46

Copper-catalyzed electrosyntheses of cyclopropanes.

of the mechanism: an ionic pathway in the noncatalyzed reaction [101] vs. a pathway involving the formation of a Ni–carbene in the nickel-catalyzed reaction [100]. The electrogenerated Ni0 activates RCHX2 by oxidative addition to give RCH(X)–NiIIX whose further two-electron reduction generates a nickel–carbene [RCH=Ni] prone to react with the alkene to form the substituted cyclopropane together with the active Ni0. Electroreductive cyclopropanations of electron-deficient alkenes are also catalyzed by CuBr, in association with pyridine, as reported by Paugam et al. (Scheme 36.46) [102]. Substituted chlorocyclopropanes are formed from trichloroalkyl derivatives [103].

H. ELECTROREDUCTIVE HYDROALKYLATION OF ELECTRON-DEFICIENT ALKENES CATALYzED BY COBALT CoI in vitamin B12 or CoIII in a vitamin B12 model catalyzes the intermolecular and intramolecular electroreductive hydroalkylation of electron-deficient alkenes, as reported by Scheffold et  al. [104,105] (Scheme 36.47). A mechanism is proposed for the intramolecular version performed from cyclohexenones substituted in the α-position by an alkyl chain with a terminal C–Br bond (R–Br), leading to bicyclic ketones (Scheme 36.47). A CoI complex activates the C–Br bond to generate a R–CoIII–Br complex (one CoIII complex has been isolated with n = 4, Scheme 36.47). Its electrochemical reduction performed at the potential of its second reduction peak generates the decalone after intramolecular nucleophilic attack of the activated C═C bond by the electrogenerated R− anion, followed by protonation [104]. Vitamin B12, cat

EWG + H+ + 2e

RX +

R EWG + X–

X = halide R = alkyl, vinyl, acyl O

O Br

(CH2)n

+ + 2e

+H

Vitamin B12 or derivative 1–20 mol%

(CH2)n

DMF, 20°C cathode Hg// Pt anode

1 mmol

n = 4, 95% n = 4, 70%

O

via the 2e reduction of n=4 Coııı Br

SCHEME 36.47

Cobalt-catalyzed electroreductive hydroalkylation of electron-deficient alkenes.

© 2016 by Taylor & Francis Group, LLC

1418

Organic Electrochemistry

IV. ELECTROCHEMICAL RECyCLINg OF THE CATALyST IN TRANSITION METAL–CATALyZED ORgANIC ELECTROSyNTHESES Some transition metal–catalyzed reactions may involve organometallic species (generated by chemical activation of an organic substrate by the transition metal) which do not require any subsequent activation by electron transfer. However, the catalytic cycle may generate in the last step of the catalytic cycle an inactive catalyst that must be reduced or oxidized back to the active form involved in an early step of the catalytic cycle. A stoichiometric amount of electrons is thus required to achieve a catalytic electrosynthesis by recycling the active catalyst at each catalytic cycle. This can be achieved by a direct or a mediated electrochemical process.

A.

DIRECT ELECTROCHEMICAL RECYCLING OF THE CATALYST

1. Electroreductive Hydrocarboxylation of Alkynes Catalyzed by Nickel The electroreductive hydrocarboxylation of alkynes catalyzed by Ni(bpy)3(BF4)2 has been reported by Duñach et al. [106–109]. The reactions, performed in an undivided cell equipped with a sacrificial Mg anode, afford a mixture of two α,β-unsaturated carboxylic acids (Scheme 36.48). The major branched α,β-unsaturated acids are obtained from terminal alkynes [106,107]. A mechanism is proposed (Scheme 36.49) [109]. The bielectronic reduction of the NiII precursor generates a Ni0 complex ligated by both CO2 and the alkyne in two different geometries. R΄

R 3 mmol

+ CO2 + 2H+ + 2e 1 atm

Ni(bpy)32+ 10 mol%

R

DMF, rt cathode C, Mg anode

HO2C

CO2H

Overall yield 60–90% Regioselectivity 65–90%

Nickel-catalyzed electroreductive hydrocarboxylation of alkynes. R΄

R R΄

R΄

R +

R = alkyl, aryl R΄ = H, alkyl

SCHEME 36.48

R΄

CO2–)2Mg2+

R

2e

Start

bpy

+ 2 bpy

CO2 2Mg2+ –)

Niıı(bpy)32+

Mg2+,H+

Ni0(bpy)2 + bpy

R΄ (bpy)Niıı

R

O

R O + (bpy)Niıı

– bpy

R΄ Ni0(bpy) O

O

R΄

R

R (bpy)Ni0

CO2 R΄ (bpy)Ni0

R΄

+ CO2

R CO2

SCHEME 36.49 Mechanism of the nickel-catalyzed electroreductive hydrocarboxylation of alkynes.

© 2016 by Taylor & Francis Group, LLC

1419

Organometallic Compounds as Tools in Organic Electrosynthesis

An intramolecular oxidative coupling between the ligated CO2 and the alkyne gives two isomeric cyclic NiII carboxylates. The cleavage of the Ni–C bond either by protons followed by NiII exchange by Mg2+ cations (generated at a sacrificial Mg anode) or by NiII exchange by Mg2+ cations followed by hydrolysis releases two α,β-unsaturated carboxylic acids and the Ni(bpy)32+ precursor. The latter is recycled back to the active Ni0 complex by a reduction at the cathode. Therefore, the activation of intermediate organometallic species by electron transfer is not required. The electrolyses are performed at the reduction potential of the Ni II precursor [109]. The electrons are only used to generate (at the very beginning of the electrolysis) and regenerate the active Ni0 complex at each catalytic cycle. The same group has developed the Ni-catalyzed electroreductive hydrocarboxylation of 1. Nonconjugated diynes (cat: Ni(bpy)3(BF4)2 or NiBr2(PMDTA), PMDTA = pentamethyldiethylenetriamine; PCO2 = 1 atm) [110] 2. 1,3-Diynes (cat: NiBr2(PMDTA); PCO2 = 1 atm) [111] 3. 1,3-Enynes (cat: NiBr2(PMDTA) or Ni(bpy)3(BF4)2; PCO2 = 1 atm) [112] 4. Alkenes (cat: NiBr2(PMDTA), PCO2 = 1 atm) [113,114] 5. Allenes (cat: NiBr2(PMDTA) or Ni(bpy)3(BF4)2; PCO2 = 1–5 atm) [115] Terminal alkynes are more reactive than internal alkynes and alkenes. 2. Electroreductive Hydroarylation of Alkenes Catalyzed by Nickel or Cobalt Nickel(II) salts catalyze the electroreductive hydroarylation of electron-deficient alkenes CH2=CH–EWG (EWG=CO2R [116], COR [116,117], CN [117]) (Scheme 36.50), as reported by Condon and Nédélec [118]. The electrolyses are carried out in an undivided cell equipped with a sacrificial Fe anode. A mechanism is proposed by the same group (Scheme 36.51, the pyridine ligands are voluntarily omitted) [117]. The bielectrochemical reduction of the NiIIBr2 salt or pyridine complex at the very beginning of the electrolysis generates a Ni0 complex stabilized by the pyridine and the activated alkene. Its oxidative addition to ArX gives Ar–NiIIX(η2-alkene). Intramolecular syn addition would generate ArCH2–CH(EWG)–NiIIX. Hydrolysis of the latter assisted by Fe2+ cations generated at the sacrificial Fe anode releases the saturated final product ArCH2–CH2–EWG together with a NiII moiety, which is reduced back to the active Ni0 at the cathode. The electrolyses are conducted at the reduction potential of the NiII precursor and no organonickel intermediate is described as being activated by electron transfer. Condon et al. have extended their reaction to heteroaryl halides [119] and vinylic halides [120]. The reactions from aryl halides have also been performed in protic solvents, such as ethanol (cat: NiBr2 and 2,2′-dipyridylamine (1:1); [41]) and in ionic liquids without any pyridine since the methylimidazolium may serve as ligand for NiII and Ni0 centers [36]. The NiBr2-catalyzed electroreductive hydroarylation of electron-deficient alkenes has been successfully used as the first step in the multistep synthesis of aromatic lactones [121,122] or benzannulated compounds [123]. The intramolecular version involving aryl halides tethered to an acrylic or fumaric moiety through an ester or amide function has been developed by Nédélec et al. [124]. ArX + 15 mmol

EWG 30 mmol

+ 2e

NiBr2 5 mol% H2O, DMF, pyridine, 60°C–80°C cathode Ni, Fe anode

Ar

– EWG + X

70–90%

X = Br, Cl Ar = m– or p-Z–C6H4 (Z = CN, COMe, CO2Et, H, OH, Me, OMe) EWG = CO2Me, CO2Et, COMe, CN

SCHEME 36.50

Nickel-catalyzed electroreductive hydroarylation of alkenes.

© 2016 by Taylor & Francis Group, LLC

1420

Organic Electrochemistry 2e

Start Ar

NiIIX2

X–

EWG

Ni0 Fe2+, H+

EWG

NiIIX EWG

Ar EWG

0

Ni

Syn addition EWG

Oxidative addition

NiII Ar

SCHEME 36.51

ArX + 7.5 mmol

ArX

X

Mechanism of the nickel-catalyzed electroreductive hydroarylation of alkenes.

COMe + H2O + 2e 10 mmol

CoBr2 13 mol%

Ar

acetonitrile/pyridine or bpy 26 mol% in DMF/pyridine cathode Ni, Fe anode

COMe + X– 22–56%

X = I, Br, Cl Ar = p-Z–C6H4 (Z = CN, COMe, CO2Me, CF3, H, OMe)

SCHEME 36.52

Cobalt-catalyzed electroreductive hydroarylation of alkenes.

In similar experimental conditions as those reported in Scheme 36.50, NiBr2 catalyzes the electroreductive arylation of electron-rich alkenes such as acrolein diethyl acetal CH2=CH–CH(OEt)2, leading to (Z)- and (E)-enol ethers ArCH2CH=CH~OEt after cleavage of an EtO − group assisted by Fe2+ electrogenerated at a Fe anode. The β-arylated aldehydes ArCH2CH2CHO are formed after hydrolysis [118,125]. Hydroarylation of electron-deficient alkenes catalyzed by CoBr2 associated with the bpy ligand has been reported by Gosmini et al. [126,127] (Scheme 36.52). However, the arylated alkene may be the major product in competition with the saturated product in the case of acrylates [127]. 3. Electrooxidative Carbonylation of Amines Catalyzed by Palladium Palladium salts or complexes catalyze the oxidative carbonylation of aliphatic and aromatic amines leading to N,N′-substituted ureas, as reported by Chiarotto and Feroci (Scheme 36.53) [128]. The reactions are performed under mild conditions at atmospheric pressure of CO and room

RNH2

+

CO

0.5 mmol 1 atm

Pd(OAc)2 10 mol% or Pd(PPh3)4 10 mol% NaOAc 2 mmol CH3CN, rt cathode Pt// C anode

RNHCONHR + 2H+ + 2e 30–83%

R = p-Z–C6H4 (Z = OMe, H, Cl); 2-pyridyl; benzyl, butyl, pentyl, cyclohexyl

SCHEME 36.53

Palladium-catalyzed electrooxidative carbonylation of amines.

© 2016 by Taylor & Francis Group, LLC

1421

Organometallic Compounds as Tools in Organic Electrosynthesis

R2 R1

R3 R4

HO

+ CO

NH2 0.5 mmol

R2

Pd(OAc)2 10 mol% NaOAc 2 mmol

R3 R4

R1

CH3CN, 50°C cathode Pt//C anode

O

+ 2H+ + 2e

NH

1 atm O 29–100%

SCHEME 36.54 Synthesis of oxazolidin-2-ones via a palladium-catalyzed electrooxidative carbonylation of β-aminoalcohols.

temperature. Sodium acetate is required to favor deprotonations. The reaction is initiated by a PdII species, Pd(OAc)2 or PdII(PPh3)22+ (generated at the anode) that is first ligated by the amine. After deprotonation of the ligated amine, CO insertion, and attack of a second amine onto the amide ligand, a Pd0 species is formed that is oxidized at a graphite anode (+0.4 V vs. SCE) to the PdII species active in the first step of the catalytic cycle. Chiarotto and Feroci have reported the electrooxidative carbonylation of 2-amino-1-alkanols to oxazolidin-2-ones catalyzed by Pd(OAc)2 (Scheme 36.54) [129]. The oxazolidin-2-one is formed via an intramolecular attack of the alcohol onto the amide ligand. The resulting Pd0 is oxidized at a graphite anode (+0.4 V vs. SCE) back to Pd(OAc)2.

B.

MEDIATED ELECTROCHEMICAL RECYCLING OF THE CATALYST

The direct recycling of the transition metal catalyst may be a problem in the absence of stabilizing ligands. A redox mediator can be used that is able to oxidize (or reduce) the short-lived transient catalyst in a fast reaction (inner-sphere process). The reduced (or oxidized) form of the mediator is then oxidized (or reduced) at an electrode at each catalytic cycle, leading to the consumption of a stoichiometric amount of electrons. 1. Electrooxidation of 1,3-Dienes Catalyzed by Palladium The electrooxidation of 1,3-cyclohexadiene in the presence of LiOAc is catalyzed by Pd(OAc)2 associated with catalytic hydroquinone, as reported by Bäckvall and Gogoll [130]. The reaction leads to cis-1,4-diacetocyclohex-2-ene or 1-aceto-4-chlorocyclohex-2-ene when performed in the presence of excess LiCl (Scheme 36.55). The catalytic cycle is initiated by the reaction of Pd(OAc)2 with the diene. After two successive nucleophilic attacks of AcO − (or AcO − and then Cl−), a Pd0 moiety is formed, which must be oxidized to the active Pd(OAc)2. Instead of using a stoichiometric amount of a chemical oxidant such as p-benzoquinone that is reduced to hydroquinone [131], a catalytic amount of the latter is used. The hydroquinone is oxidized to p-benzoquinone at a Ti anode (via catalytic MnO2 that serves as a second redox mediator) or directly at a Pt anode. This mediated electrochemical recycling of Pd0 to PdII via a chemical oxidant such as p-benzoquinone, made catalytic by the electrochemical oxidation of OAc + 2H+ + 2e

AcO + 2HOAc

Pd(OAc)2 5–10 mol% hydroquinone 25 mol% LiOAc, HOAc, rt cathode steel//Ti/MnO2 or Pt anode

72–74% LiCl Cl + 2H+ + 2e

AcO 46% Major cis

SCHEME 36.55

Palladium-catalyzed electrooxidation of 1,3-dienes.

© 2016 by Taylor & Francis Group, LLC

1422

Organic Electrochemistry Pd0 + O

O + 2H+

Pdll2+ + HO

OH

–2e, –2H+

SCHEME 36.56

Recycling of p-benzoquinone via electrooxidation of hydroquinone.

hydroquinone, is a pioneer contribution by Bäckvall et al. (Scheme 36.56) [130]. It is a very versatile procedure as will be illustrated in the following. 2. Electrooxidative Heck-Type Reactions from Arenes Catalyzed by Palladium Classical Heck reactions performed from aryl halides are catalyzed by palladium(0) complexes, but halide ions are released as waste [3,132]. Pd-catalyzed Heck-type reactions from arenes, ArH (cheap, widely spread, no waste), as pioneered by Jia and Fujiwara [133] but working under mild conditions remains a challenge (Equation 1 in Scheme 36.57). Such reactions require a palladium(II) catalyst able to activate the Ar–H bond and a stoichiometric amount of an oxidant (such as p-benzoquinone) to oxidize the Pd0 formed at the end of every catalytic cycle, back to the active PdII catalyst [134]. Amatore et al. have developed Pd(OAc)2-catalyzed electrooxidative Heck-type reactions from an arene and alkenes (n-butyl acrylate, styrene) in the presence of a catalytic amount of p-benzoquinone (or hydroquinone) (Equation 2 in Scheme 36.57) [135]. The reactions are performed in acetic acid at room temperature. Protons are reduced at a Ni cathode. Despite the use of a Ni cathode known to reduce protons at a low potential, the latter remains still more negative than the reduction potential of p-benzoquinone, preventing the use of a one-compartment cell. Most steps of the catalytic cycles are similar to those of a classical Heck reaction performed from ArX. The aryl-PdII complex prone to react with the alkene is formed by the activation of the Ar–H bond by Pd(OAc)2 (Scheme 36.58). This key step is favored when the arene has a substituent Z (the amido group in the present case) as an arm able to coordinate Pd(OAc)2 to induce an intramolecular orthopalladation (Scheme 36.58) [136]. After the classical carbopalladation, C–C internal rotation and β-hydride elimination, a reductive elimination from HPd(OAc) generates a transient Pd0 moiety [132] that is oxidized back to the active PdII by p-benzoquinone. The latter is used in a catalytic amount because it is recycled by oxidation at the anode of the hydroquinone formed in the chemical oxidation of Pd0 (Schemes 36.56 and 36.58). The catalytic cycle does not involve any activation of intermediate organopalladium complexes by electron transfer, but a stoichiometric amount of electrons is required to recycle the p-benzoquinone at every catalytic cycle. In contrast to Heck reactions involving aryl halides that are performed at high temperatures in the presence of a base (required in the reductive elimination of HX), the reactions from ArH are performed at room temperature, without any base but in acetic acid. The reversible reductive elimination of [Pd0] from HPd(OAc) is shifted by the p-benzoquinone, which acts as a ligand and a chemical oxidant of the transient Pd0 (inner-sphere process) [131].

Ar–H +

R + Oxidant

PdII cat

Ar

H

H Pd(OAc)2 10 mol% p-benzoquinone 10 mol%

N + O

R

H 4.4 mmol

SCHEME 36.57

(1)

R

4.8 mmol

HOAc, rt cathode Ni//C anode R = CO2nBu (82%) R = Ph (32%)

N O

+ 2H+ + 2e

R

Palladium-catalyzed electrooxidative heck-type reaction from arenes.

© 2016 by Taylor & Francis Group, LLC

(2)

1423

Organometallic Compounds as Tools in Organic Electrosynthesis H e

d no

N

A

–2e

hydroquinone

O

Start H

Oxidation –2H+

+2H+

AcO–

Pd(OAc)2

HOAc

p-benzonquinone 10 mol%

C–H activation H N

[Pd0]

O PdII

HOAc

1/2 AcO

Reductive elimination

2

Carbopalladation H

Pd

OAc R OAc

H

H

β-hydride elimination

N O

Pd

Z Ar

OAc Pd (Z = NHCOMe)

Ar H

Z

H R

H H

H R

C–C internal rotation R

SCHEME 36.58

Mechanism of a palladium-catalyzed electrooxidative heck-type reaction from arenes.

3. Electrooxidation of Alcohols Catalyzed by Palladium Amatore et  al. have reported the electrochemical oxidation of primary or secondary alcohols to aldehydes or ketones, respectively, under anaerobic conditions, in the presence of a base and a catalytic amount of Pd(OAc)2 associated with catalytic p-benzoquinone (Scheme 36.59) [137]. A mechanism is proposed in Scheme 36.60 [137]. The investigated secondary or primary alcohols do not exhibit any oxidation peak in DMF (potentials up to +2 V vs. SCE at a glassy carbon disk electrode). No oxidation peak is observed after addition of K2CO3 as a base, which suggests a slow deprotonation of the alcohol at 25°C. At high temperature (80°C), a PdII-alkoxide is formed by substitution of one acetate ligand of Pd(OAc)2 by the alkoxide. It undergoes a β-hydride elimination that generates the aldehyde (or the ketone) and HPd(OAc). A reductive elimination affords a Pd0 moiety that is oxidized to the active PdII catalyst by p-benzoquinone. The latter is made catalytic by oxidation of hydroquinone at the anode (Scheme 36.60). Pd(OAc)2 and Pd0(PPh3)4 may be used as catalysts without any p-benzoquinone via the direct oxidation of the Pd0 moiety at the anode, but the reactions are slower and cannot be made to completion due to catalyst decomposition [137].

RCH(OH)R΄ 2 mmol

Pd(OAc)2 10 mol% p-benzoquinone 10 mol% K2CO3, DMF, 80°C cathode Ni // C anode

RCOR΄ + 2H+ + 2e 41–98%

R = aryl, alkenyl, alkyl R΄= H, aryl, alkyl

SCHEME 36.59

Palladium-catalyzed electrooxidation of alcohols.

© 2016 by Taylor & Francis Group, LLC

1424

Organic Electrochemistry e od An –2e

RR΄CH-OH hydroquinone

Oxidation –2H+

Start

Base

Pd(OAc)2

+2H+

p-benzoquinone 10 mol%

RR΄CH-O–

AcO– AcO–

[Pd0]

AcO

HOAc

H R

O

Reductive elimination AcO

Pd

R΄ Pd

H β-hydride elimination

RCOR΄

SCHEME 36.60 Mechanism of the palladium-catalyzed electrooxidation of alcohols.

This anaerobic electrochemical procedure is quite general and selective and avoids the formation of H2O2 generated in the aerobic Pd(OAc)2-catalyzed oxidation of alcohols [138] whose rate is moreover limited by the rate-determining slow dissolution and low solubility of O2 into most organic solvents [139]. 4. Electrooxidative Homocoupling of Arylboron Derivatives Catalyzed by Palladium The homocoupling of arylboronic acids or arylboronates is catalyzed by palladium complexes ligated by a phosphine ligand. It requires a stoichiometric amount of an oxidant, which can be air or dioxygen [140–142]. A new anaerobic electrooxidative homocoupling of arylboronic acids, arylboronates, and aryltrifluoroborates has been developed in the presence of catalytic amounts of ligandless Pd(OAc)2 and catalytic p-benzoquinone, which serves as a redox mediator for the oxidation of the transient Pd0 to the active PdII species (Scheme 36.61) [143]. The electrooxidative homocouplings are performed in DMF, DMF/water mixture, or pure water (for ArBF3− with the cheap and water-soluble supporting electrolyte Na2SO4). This procedure avoids the use of stoichiometric chemical oxidants (as O2), the formation of peroxopalladium moieties, and subsequent by-products (phenols) [142]. The mechanism of the Pd-catalyzed electroreductive homocoupling of arylboronic acids ArB(OH)2 is proposed in Scheme 36.62 [143]. A double transmetallation of Pd(OAc)2 by the

–

2 Ar–B(OR΄)2 or ArBF3 1 mmol

Pd(OAc)2 10 mol% p-benzoquinone 10 mol% Ar–Ar + 2e DMF, H2O, 80°C cathode Ni // C anode

32–99%

R΄= alkyl; Ar = p-Z–C6H4 (Z = CN, H) R΄= H, Ar = p-Z–C6H4 (Z = CN, Br, F, H, Me, OMe); naphthyl

SCHEME 36.61

Palladium-catalyzed electrooxidative homocoupling of arylboron derivatives.

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1425

Organometallic Compounds as Tools in Organic Electrosynthesis e

Oxidation

–2e

hydroquinone

od

An

–2H+

2 ArB(OH)2

Pd(OAc)2

2H+ p-benzoquinone 10 mol%

Start

Transmetallation 2AcO–

2AcO– [ Pd0]

Ar–Pd–Ar Reductive elimination

Ar–Ar

SCHEME 36.62 Mechanism of the palladium-catalyzed electrooxidative homocoupling of arylboronic acids. Pd(OAc)2 10 mol% TEMPO 30 mol% 2 Ar-B(OR’)2 Acetonitrile, H2O, rt, K2CO3 0.2 mmol cathode Pt // Pt anode

Ar–Ar + 2e 23–97%

R΄= alkyl; Ar = p-Z–C6H4 (Z = NO2, Cl, H, Me, OMe, NMe2) R΄= H; Ar = p-Z–C6H4 (Z = NO2, CO2Et, Cl, H, Me, tBu, OMe, OPh); naphthyl

SCHEME 36.63

Palladium-catalyzed electrooxidative homocoupling of arylboronates.

arylboronic acid generates a bisarylpalladium(II) intermediate. A reductive elimination gives the biaryl and the transient [Pd0] that is converted to Pd(OAc)2 by oxidation by the p-benzoquinone made catalytic by oxidation of the hydroquinone at the anode at each catalytic cycle. Tanaka et al. have reported the use of catalytic TEMPO (2,2′-tetramethylpiperidyloxide) in the electrooxidative homocoupling of arylboronic acids and esters catalyzed by Pd(OAc)2 (Scheme 36.63) [144,145]. The oxidation of TEMPO at the anode generates TEMPO+ that oxidizes the transient [Pd0] formed at the end of each catalytic cycle to a cationic transient active [PdII(CH3CN)4]2+. TEMPO is regenerated in this chemical oxidation. 5. Electrooxidative Wacker Reactions Catalyzed by Palladium The Wacker process, oxidation of alkenes RCH=CH2 to aldehydes or ketones, is catalyzed by PdIIX2 salts (e.g., PdCl2) [146]. The mechanism (still under debate) involves a complexation of the alkene to the PdII salt, followed by the nucleophilic trans-attack of H2O onto the ligated alkene with two possible attack sites according to the nature of the R group. A β-hydride elimination from RCH(OH)– CH2–PdX or from RCH(PdX)–CH2OH delivers the aldehyde or the ketone, respectively, and HPdX whose reductive elimination gives a [Pd0] moiety. The latter must be recycled back to the active PdII by oxidation. The chemical oxidant is often a CuII salt (e.g., CuCl2) used in catalytic amount because recycled by the oxidation of the resulting CuI by dioxygen [146]. Torii et al. have reported the first electrochemical version of the Wacker process, using PdCl2 or Pd(OAc)2 as catalyst and a catalytic amount of (p-Br–C6H4)3N as redox mediator (Scheme 36.64) [147]. The iminium radical cation generated at the anode oxidizes [Pd0] back to PdII. A catalytic amount of p-benzoquinone has also been used [147].

+ H2O R 1 mmol

Pd(OAc)2 5 mol% (p-BrC6H4)3N 20 mol% or p-benzoquinone 20 mol% CH3CN, rt cathode Pt // Pt anode

O R 41–90%

SCHEME 36.64 Palladium-catalyzed electrooxidative Wacker reactions.

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+ 2H+ + 2e

1426

Organic Electrochemistry O

Pd(OAc)2 10 mol% TEMPO 30 mol% + H2O R CH3CN, rt 0.5 mmol cathode Pt // Pt anode R = Alkyl; Ph

SCHEME 36.65

; Ph

+ 2H+ + 2e

R 63–91%

O

Palladium-catalyzed electrooxidative Wacker reactions.

Tanaka et al. have reported similar reactions using Pd(OAc)2 as catalyst and a catalytic amount of TEMPO as a redox mediator for the oxidation of [Pd0] (Scheme 36.65) [148]. According to the authors, an active cationic Pd(CH3CN)42+ complex is formed via the oxidation of the acetate ligands of Pd(OAc)2 at the very beginning of the electrolysis [148]. 6. Electrooxidative Asymmetric Dihydroxylation of Alkenes Catalyzed by Osmium OsO4 associated with a chiral ligand (L* = dihydroquinidine DHQD, dihydroquinine DHQ or derivatives such as (DHQD)2PHAL [PHAL = phthalazine]) reacts with alkenes to give asymmetric dihydroxylation [149]. The reactions are made catalytic in osmium in the presence of a chemical oxidant used in stoichiometric amount that oxidizes the OsVI species formed in the catalytic cycle back to the active OsVIII species: for example, N-methylmorpholine N-oxide (NMO) associated with catalytic OsO4 [149], or K3Fe(CN)6 [149] associated with catalytic K2OsVIIIO2(OH)4 in tBuOH/H2O (Scheme 36.66). An electrochemical oxidation has been developed as an alternative to regenerate OsO4 from OsVI species, either directly [150,151] or mediated by a chemical oxidant used in catalytic amount (K3Fe(CN)6) [150–152] or I2 [153] which is recycled by oxidation at the anode of its reduced form. The electrosyntheses are performed in undivided cells. The enantioselectivity of the electrochemical process is close to that obtained in the presence of a stoichiometric amount of the oxidant (Scheme 36.67). R + 2H2O + Oxidant R

OsVlll or OsVl 0.01–5 mol%, L*

R

OH

R

R

OH

O via O

O Os L*

SCHEME 36.66

O R

Osmium-catalyzed asymmetric dihydroxylation of alkenes. OH OsO4 0.1 mol% K3Fe(CN)6 40 mol% (DHDQ)2PHAL 1.5 mol%

OH + 2e

K2CO3, tBuOH/H2O 100%, 91% ee

1 mmol K2OsO2(OH)4 0.2 mol% (DHDQ)2PHAL 1 mol% K2CO3, tBuOH/H2O, 0°C Ph 1 mmol

cathode Pt, Pt anode

H HO

OH Ph

I2 50 mol%

98%, 93.7% ee

K3Fe(CN)6 10 mol%

95%, 97.3% ee

+ 2e

SCHEME 36.67 Osmium-catalyzed electrooxidative asymmetric dihydroxylation of alkenes.

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Organometallic Compounds as Tools in Organic Electrosynthesis

7. Electroreductive Carbonylation of Aryl Halides Catalyzed by Palladium Pd0(PPh3)4 or PdCl2(PPh3)2 catalyzes the electrosyntheses of aromatic aldehydes from aryl iodides and carbon monoxide in the presence of formic acid (Scheme 36.68) [153]. The electrosyntheses consume one electron per mole of aryl iodide and are all performed at the same potential (−1.5 V vs. SCE) whatever the aryl halide. The mechanism of the reaction catalyzed by Pd0(PPh3)4 has been investigated by Amatore et al. (Scheme 36.69) [154,155]. An acyl-palladium(II) complex ArCOPdI(PPh3)2 is generated by carbonylation of ArPdI(PPh3)2 formed in the oxidative addition of Pd0(PPh3)2 to ArI (Scheme 36.69). The electrolyses are performed at the same potential (whatever the aryl iodide), which is less negative than the reduction potential of ArCOPdI(PPh3)2 [156], indicating that its reduction is not required. ArCOPdI(PPh3)2 is trapped by formate ions HCO2−, which are generated in the reversible oxidative addition of Pd0(PPh3)2 to formic acid (Scheme 36.69) [155]. A classical decarboxylation in ArCOPd(OCHO)(PPh3)2 generates ArCOPdH(PPh3)2 and then the aldehyde ArCHO and Pd0(PPh3)2 by reductive elimination (Scheme 36.69). Part of the Pd0 is consumed in its reversible oxidative addition to formic acid [155], which generates the cationic hydrido complex, HPdII(DMF)(PPh3)2+, evidenced by conductivity measurement (Scheme 36.69) [156,157]. The latter is reduced at the cath-

Arl + CO + HCO2H + 1e 1 mmol 0.5 mmol 1atm

PdCl2(PPh3)2 10 mol% or Pd0(PPh3)4 10 mol%

ArCHO + CO2 + 1/2 H2 + l– 40–93 %

DMF, 60°C cathode C // Pt anode

Ar = p-Z-C6H4 (Z = Me, OH, H) o-Me–C6H4 ; m,o- Me2–C6H3 ; m,p-Me2–C6H3

SCHEME 36.68 Palladium-catalyzed electroreductive carbonylation of aryl iodides.

de ho HPdllSL2+ t Ca Reduction

DMF L = PPh3 K = 10–3 M (0°C)

K

1e

HCO2H

1/2 H2

Pd0L2 –L

ArCHO

Pd0L3

Reductive elimination

HCO2–

Arl CO Oxidative addition

Pd0L2(CO)

–L Pd0L4

ArCOPdllHL2

ArPdllIL2

Start

CO CO2 ArCOPdll(OCHO)L2

l–

ArCOPdllIL2

HCO2–

SCHEME 36.69 Mechanism of the palladium-catalyzed electroreductive carbonylation of aryl iodides.

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1428

Organic Electrochemistry

Start

Oct

+

+

N

N

C8V2+, 1 mmol 2e

(1)

Oct

(2)

(3) Recycling after work-up

THF , rt cathode Pt, Mg anode

N

C8V0

Oct

N

Oct

Transfer

ArBr + C8V0 0.5 mmol 1 mmol

PdCl2(PhCN)2 5mol% THF, 60°C

1/2 ArAr + Br– + 1/2 C8V2+ 71–96%

Ar = p-Z–C6H4 (Z = MeO, Me, OH, H, Cl, CO2Me, CN)

SCHEME 36.70 quinoid.

Palladium-catalyzed electroreductive homocoupling of aryl bromides mediated by

ode, back to the active Pd0(PPh3)2 [154]. In such reactions, the Pd0 plays a double role, activation of the aryl iodide and formic acid. Interestingly, the Pd0 complex catalyzes the reduction of the protons of acetic acid to H2 via the reduction of HPdII(DMF)(PPh3)2+: Ep = −0.99 V instead of Ep = −2.0 V vs. SCE for acetic acid, at a glassy carbon disk electrode in DMF [155]. 8. Palladium-Catalyzed Homocoupling of Aryl Halides Mediated by Electrogenerated Quinoid C8V0 The electroreductive homocoupling of aryl halides to biaryls catalyzed by PdII complexes required two reduction steps at a cathode: (1) reduction of PdII to Pd0 and (2) reduction of an ArPdX complex (Schemes 36.38 and 36.39). Tanaka et al. have reported an alternative procedure in which the reaction starts by the bielectronic reduction of a stoichiometric amount of N,N′-dioctyl-4,4′bipyridinium (C8V2+,Tf 2N−) to the quinoid C8V0 in an undivided cell (Scheme 36.70) [158]. The blue solution of C8V0 is then transferred to a THF solution containing the aryl halide ArX and the PdII catalyst. Chemical homocoupling takes place since C8V0 is a strong reducing agent that reduces both the PdII and the ArPdX complex, leading to the synthesis of the biaryl (Scheme 36.70). Once the chemical reaction is over, C8V2+ formed in the chemical oxidation of C8V0 can be reduced again at a cathode and the ensuing C8V0 recycled into a new batch of ArX and catalyst (Scheme 36.70).

V. CONCLUSION Transition metal complexes are efficient catalysts for organic electrosyntheses. The investigation of their mechanisms by means of electrochemical techniques reveals the intricate role of chemical activation steps by transition metals and activation steps by electron transfers. By using a transition metal catalyst for organic electrosyntheses, it is possible (1) to selectively activate bonds that would not be activated by simple electron transfer and consequently to trigger a reaction toward a direction opposite to that observed in the absence of catalyst and (2) to avoid the formation of organic radicals that might be formed in noncatalyzed electrosyntheses, as described in many chapters of this monograph and in a review chapter [159]. This explains why most reactions are stereo- or regiospecific.

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Organometallic Compounds as Tools in Organic Electrosynthesis

1429

By finely tuning the potential of an electrode, it is possible: (1) to generate very active transition metal catalysts from stable and unreactive precursors at the very beginning of the electrosyntheses; (2) to activate intermediate organometallic species formed in chemical activation steps by one- or two-electron transfer according to the nature of the metal and ligand (as a consequence, complexes with unusual oxidation states [NiI, NiIII] or unusual anionic species [Pd0, PdII] are formed); (3) to activate organometallic species that become more easily reduced than the initial substrates, at the origin of the catalysis; (4) to invert the reactivity of organic substrates toward electrophilic reactions instead of nucleophilic reactions usually observed in the absence of an electron source; (5) to avoid the use of stoichiometric chemical oxidants or reductants by the electrochemical recycling of active catalysts at each catalytic cycle; and (6) to investigate the mechanism of transition metal–catalyzed reactions by the characterization of electrochemical steps (potential and number of electron(s)) and the characterization of the kinetics of chemical steps that do not involve any electron transfer (determination of the rate constants).

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Organometallic Compounds as Tools in Organic Electrosynthesis 79. 80. 81. 82. 83. 84. 85. 86. 87. 88. 89. 90. 91. 92 93. 94. 95. 96. 97. 98. 99. 100. 101. 102. 103. 104. 105. 106. 107. 108. 109. 110. 111. 112. 113. 114. 115. 116. 117. 118. 119. 120. 121. 122. 123. 124. 125. 126. 127. 128. 129. 130.

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Jutand, A.; Négri, S. Eur. J. Org. Chem. 1998, 1811–1821. Jutand, A.; Négri, S.; Mosleh, A. J. Chem. Soc. Chem. Commun. 1992, 1729–1730. Brunel, J.; Jutand, A.; Ledoux, I.; Zyss, J.; Blanchard-Desce, M. Synth. Met. 2001, 124, 195–199. Jutand, A.; Mosleh, A. Organometallics 1995, 14, 1810–1817. Amatore, C.; Carré, C.; Jutand, A. Acta Chem. Scand. 1998, 52, 100–106. Lindsey, A. S.; Jesse, H. Chem. Rev. 1957, 57, 583–620. Jutand, A.; Négri, S. Synlett 1997, 719–721. Kamekawa, H.; Senboku, H.; Tokuda, M. Tetrahedron Lett. 1998, 39, 1591–1594. Jutand, A.; Negri, S. Organometallics 2003, 22, 4229–4237. Senboku, H.; Fujimura, Y.; Kamekawa, H.; Tokuda, M. Electrochim. Acta 2000, 45, 2995–3003. Kamekawa, H.; Senboku, H.; Tokuda, M. Electrochim. Acta 1997, 42, 2117–2123. Kamekawa, H.; Kudoh, H.; Senboku, H.; Tokuda, M. Chem. Lett. 1997, 917–918. Torii, S.; Tanaka, H.; Katoh, T.; Morisaki, K. Tetrahedron Lett. 1984, 33, 3207–3208. Medeiros, M. J.; Pintaric, C.; Duñach, E. Electrochim. Acta 2011, 56, 4384–4389. Torii, S.; Tanaka, H.; Morisaki, K. Tetrahedron Lett. 1985, 26, 1655–1658. Amatore, C.; Carré, E.; Jutand, A.; Tanaka, H.; Ren, Q.; Torii, S. Chem. Eur. J. 1996, 2, 957–966. Jutand, A.; Mosleh, A. J. Org. Chem. 1997, 62, 261–274. Fitton, P.; Rick, E. A. J. Organomet. Chem. 1971, 28, 287–288. Fauvarque, J. F.; Pflüger, F.; Troupel, M. J. Organomet. Chem. 1981, 208, 419–427. Marzouk, H.; Rollin, Y.; Folest, J. C.; Nédélec, J.-Y.; Périchon, J. J. Organomet. Chem. 1989, 369, C47–C50. Amatore, C.; Jutand, A.; Périchon, J.; Rollin, Y. Monatsh. Chem. 2000, 131, 1293–1304. Sengmany, S.; Léonel, E.; Paugam, J.-P.; Nédélec, J.-Y. Tetrahedron 2002, 58, 271–277. Léonel, E.; Paugam, J.-P.; Condon-Gueugnot, S.; Nédélec, J.-Y. Tetrahedron 1998, 54, 3207–3212. Oudeyer, S.; Léonel, E.; Paugam, J.-P.; Nédélec, J.-Y. Tetrahedron 2003, 59, 1073–1081. Oudeyer, S.; Léonel, E.; Paugam, J.-P.; Sulpice-Gaillet, C.; Nédélec, J.-Y. Tetrahedron 2006, 62, 1583–1589. Scheffold, R.; Dike, M.; Dike, S.; Herold, T.; Walder, L. J. Am. Chem. Soc. 1980, 102, 3642–3644. Scheffold, R.; Abrecht, S.; Orfinski, R.; Ruf, H.-R.; Stamouli, P.; Tinembart, O.; Walder, L.; Weymuth, C. Pure App. Chem. 1987, 59, 363–372. Duñach, E.; Périchon, J. J. Organomet. Chem. 1988, 352, 239–246. Labbé, E.; Duñach, E.; Périchon, J. J. Organomet. Chem. 1988, 353, C51–C56. Duñach, E.; Derien, S.; Périchon, J. J. Organomet. Chem. 1989, 364, C33–C36. Derien, S.; Duñach, E.; Périchon, J. J. Am. Chem. Soc. 1991, 113, 8447–2454. Derien, S.; Duñach, E.; Périchon, J. J. Organomet. Chem. 1990, 385, C43–C46. Derien, S.; Clinet, J.-C.; Duñach, E.; Périchon, J. J. Chem. Soc. Chem. Commun. 1991, 549–550. Derien, S.; Clinet, J.-C.; Duñach, E.; Périchon, J. J. Organomet. Chem. 1992, 424, 213–224. Derien, S.; Clinet, J.-C.; Duñach, E.; Périchon, J. Tetrahedron 1992, 48, 5235–5248. Derien, S.; Clinet, J.-C.; Duñach, E.; Périchon, J. J. Org. Chem. 1993, 58, 2578–2588. Derien, S.; Clinet, J.-C.; Duñach, E.; Périchon, J. Synlett 1990, 361–364. Condon-Gueugnot, S.; Léonel, E.; Nédélec, J.-Y.; Périchon, J. J. Org. Chem. 1995, 60, 7684–7686. Condon, S.; Nédélec, J.-Y.; Périchon, J. Eur. J. Org. Chem. 2002, 105–111. Condon, S.; Nédélec, J.-Y. Synthesis 2004, 3070–3078. Condon, S.; Dupré, D.; Lachaise, I.; Nédélec, J.-Y. Synthesis 2002, 1752–1758. Condon-Gueugnot, S.; Dupré, D.; Nédélec, J.-Y.; Périchon, J. Synthesis 1997, 1457–1460. Métay, E.; Léonel, E.; Sulpice-Gaillet.; Nédélec, J.-Y. Synthesis 2005, 1682–1688. Métay, E.; Léonel, E.; Condon, S.; Nédélec, J.-Y. Tetrahedron 2006, 62, 8515–8524. Condon, S.; El Ouarradi, A.; Métay, E.; Léonel, E.; Bourdonneau, M.; Nédélec, J.-Y. Tetrahedron 2008, 64, 9388–9395. de Mendonça Cavalcanti, J. C.; Goulart, M. O. F.; Léonel, E.; Nédélec, J.-Y. Tetrahedron Lett. 2002, 43, 6343–6345. Condon, S.; Dupré, D.; Nédélec, J.-Y. Org. Lett. 2003, 5, 4701–4703. Gomes, P.; Gosmini, C.; Nédélec, J.-Y.; Périchon, J. Tetrahedron Lett. 2000, 41, 3385–3388. Gomes, P.; Gosmini, C.; Nédélec, J.-Y.; Périchon, J. Tetrahedron Lett. 2002, 43, 5901–5903. Chiarotto, I.; Feroci, M. J. Org. Chem. 2003, 68, 7137–7139. Chiarotto, I.; Feroci, M. Tetrahedron Lett. 2001, 42, 3451–3453. Bäckvall, J.-E.; Gogoll, A. J. Chem. Soc. Chem. Commun. 1987, 1237–1238.

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131. 132. 133. 134.

Grennberg, H.; Gogoll, A.; Bäckvall, J.-E. Organometallics 1993, 12, 1790–1793. Heck, R. F. J. Am. Chem. Soc. 1969, 91, 6707–6714. Jia, C.; Fujiwara, Y. Pure Appl. Chem. 2001, 73, 319–324. Boele, M. D. K.; van Strijdonck, G. P. F.; de Vries, A. H. M.; Kamer, P. C. J.; de Vries, J. G.; van Leeuwen, P. W. N. M. J. Am. Chem. Soc. 2002, 124, 1586–1587. Amatore, C.; Cammoun, C.; Jutand, A. Adv. Synth. Catal. 2007, 349, 292–296. For a review on orthometallation, see: Pfeffer, M. Pure Appl. Chem. 1992, 64, 335–342. Amatore, C.; Cammoun, C.; Jutand, A. Synlett 2007, 2173–2178. Peterson, K. P.; Larock, R. C. J. Org. Chem. 1998, 63, 3185–3189. Steinhoff, B. A.; Stahl, S. S. J. Am. Chem. Soc. 2006, 128, 4348–4355. Moreno-Mañas, M.; Pérez, M.; Pleixats, R. J. Org. Chem. 1996, 61, 2346–2351. Yoshida, H.; Yamaryo, T.; Ohshita, J.; Kunai, A. Tetrahedron Lett. 2003, 44, 1541–1544. Adamo, C.; Amatore, C.; Ciofini, I.; Lakmini, H.; Jutand, A. J. Am. Chem. Soc. 2006, 128, 6829–6836. Amatore, C.; Cammoun, C.; Jutand, A. Eur. J. Org. Chem. 2008, 4567–4570. Mitsudo, K.; Shiraga, T.; Tanaka, H. Tetrahedron Lett. 2008, 49, 6593–6395. Mitsudo, K.; Shiraga, T.; Kagen, D.; Shi, D.; Becker, J. Y.; Tanaka, H. Tetrahedron 2009, 65, 8384–8388. Smidt, V. J.; Hafner, W.; Jita, R.; Sedlmeier, J.; Sieber, J.; Rüttinger, R.; Kojer, H. Angew. Chem. 1959, 71, 176–182. Inokuchi, T.; Ping, L.; Hamaue, F.; Izawa, M.; Torii, S. Chem. Lett. 1994, 121–124. Mitsudo, K.; Kaide, T.; Nakamoto, E.; Yoshida, K.; Tanaka, H. J. Am. Chem. Soc. 2007, 129, 2246–2247. Kolb, H. C.; Van Nieuwenhze, M. S.; Sharpless, K. B. Chem. Rev. 1994, 94, 2483–2547. Gao, Y.; Zepp, C. M.; Wai, J. S. M. Abstracts of papers of the Am. Chem. Soc. 1990, 199, 15–15. Gao, Y.; Zepp, C. M.; 1993, WO Pat. 9 317 150. Torii, S.; Liu, P.; Tanaka, H. Chem. Lett. 1995, 319–320. Torii, S.; Liu, P.; Bhuvaneswari, N.; Amatore, C.; Jutand, A. J. Org. Chem. 1996, 61, 3055–3560. Carelli, I.; Chiarotto, I.; Cacchi, S.; Pace, P.; Amatore, C.; Jutand, A.; Meyer, G. Eur. J. Org. Chem. 1999, 1471–1473. Amatore, C.; Jutand, A.; Meyer, G.; Carelli, I.; Chiarotto, I. Eur. J. Inorg. Chem. 2000, 1855–1859. Amatore, C.; Carré, E.; Jutand, A.; Tanaka, H.; Torii, S.; Chiarotto, I.; Carelli, I. Electrochim. Acta 1997, 42, 2143–2152. Jutand, A. Eur. J. Inorg. Chem. 2003, 2017–2040. Kuroboshi, M.; Kobayashi, R.; Nakagawa, T.; Tanaka, H. Synlett 2009, 85–88. Schäfer, H. J. In: Organic Electrochemistry in Encyclopedia of Electrochemistry, Bard, A. J.; Stratmann, M. (Eds.), Wiley-VCH, Weinheim, Germany, 2004.

135. 136. 137. 138. 139. 140. 141. 142. 143. 144. 145. 146. 147. 148. 149. 150. 151. 152. 153. 154. 155. 156. 157. 158. 159.

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Section VI Stereochemical and Biological Aspects

© 2016 by Taylor & Francis Group, LLC

37

Electrosynthesis of Bioactive Materials Randi K. Gbur and R. Daniel Little

CONTENTS I. II.

Introduction ........................................................................................................................ 1436 N-Oxidation ....................................................................................................................... 1436 A. Alkaloids .................................................................................................................... 1436 1. Levobupivacaine .................................................................................................. 1436 2. Tetrahydroisoquinoline Alkaloids ....................................................................... 1437 III. Coupling of Electron-Rich Aromatics ............................................................................... 1438 A. O-Methylthalibrine ..................................................................................................... 1438 IV. Electrogenerated Reagents ................................................................................................. 1439 A. Hypervalent Iodine ..................................................................................................... 1439 1. Tetrahydropyrroloiminoquinone alkaloids .......................................................... 1439 B. Methoxy Radical ........................................................................................................ 1439 1. En Route to Parasitenone..................................................................................... 1439 V. Umpolung ........................................................................................................................... 1441 A. Bicyclo [3.2.1] Skeleton .............................................................................................. 1441 B. (−)-Crobarbatic Acid, (+)-Nemorensic Acid ............................................................... 1441 VI. Terpenoids .......................................................................................................................... 1443 A. Furan-Olefin Coupling ............................................................................................... 1443 1. Eunicellin Diterpenes .......................................................................................... 1443 2. Cyathin Core........................................................................................................ 1443 3. Tricholomalide A ................................................................................................. 1444 4. (−)-Alliacol A ...................................................................................................... 1446 5. Arteannuin Skeleton ............................................................................................ 1448 6. (−)-Heptemerone B and (−)-Guanacastepene E................................................... 1449 VII. Mediators ........................................................................................................................... 1450 A. Daucene ...................................................................................................................... 1450 VIII. Oxidative Dearomatization .................................................................................................1451 A. Introductory Remarks .................................................................................................1451 B. Heliannuol E ............................................................................................................... 1453 IX. Cycloaddition Reactions .................................................................................................... 1454 A. Quinone Methides ...................................................................................................... 1454 1. Euglobals ............................................................................................................. 1454 B. Azaquinones ............................................................................................................... 1454 1. Neuroprotective Agents ....................................................................................... 1454

1435

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1436

Organic Electrochemistry

X.

Carbohydrate Analogues and Precursors........................................................................... 1455 A. Reductive Dehalogenation: Inositol and Conduritol Synthons .................................. 1455 B. Glycals ........................................................................................................................ 1456 C. Valienamine Analogues ............................................................................................. 1457 XI. Closing Remarks ................................................................................................................ 1458 Acknowledgments........................................................................................................................ 1458 References .................................................................................................................................... 1458

I. INTRODUCTION The tremendous structural diversity of natural products, coupled with their intriguing and oftentimes useful bioactivities, continues to interest and stimulate the creativity of chemists. We note that the role of electrochemistry in the design and synthesis of these materials is increasing, just as the number of persons who realize the importance and utility of electrochemical tools is increasing. It seems clear that the important advances that continue to be made by electrochemists are opening the eyes of traditional organic chemists who have previously not embraced the field [1–7]. They have begun to more fully appreciate that redox chemistry provides many important tools that can be used to make new bonds and construct complex frameworks, and that it is much more than simply a means to oxidize an alcohol or reduce a carbonyl unit. A greater appreciation for the fact that many electrochemical processes accord with the tenants of green chemistry and provide methods that contribute to sustainability is also important [8,9]. This chapter summarizes some of the progress that has been made since the fourth edition of this book was published in 2001. It is not intended to be all encompassing in its scope and concludes with literature coverage through 2011. We apologize to colleagues whose important research we have not discussed; oversight is unintended. Also, the chapter does not repeat the chemistry we described in Chapter 19 of the fourth edition [10]. The reader who is interested in the synthetic utility of the Kolbe oxidation, for example, is referred to the fourth edition.

II.

N-OXIDATION [11]

A.

ALKALOIDS

1. Levobupivacaine Bupivacaine, the most widely used local anesthetic during childbirth in the United States and Europe, has serious neurotoxic side effects for both mother and child. Recently, less toxic alternatives, viz., levobupivacaine (1) and ropivacaine (2), have shown promising results in clinical studies [12]. Kumar and Ramachandran disclosed an asymmetric synthesis of levobupivacaine (1) in 2005 [13], but very few methods exist for an efficient enantioselective generation of the pipecolic acid moiety. In 2008, Shankaraiah, Pilli, and Santos took advantage of the selective oxidation of cyclic N-carbamates under cation pool conditions in order to form the stereogenic center [14,15]. Thus, electrolysis of 4 at 0°C for 5 h, using a platinum anode and tungsten cathode, generated the pooled N-acyliminium cation, 5, that was intercepted in situ by cyanide. O N

–2e, –H

low T OR* 4 R* = 8-phenylmenthyl

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–

+ +

O

N

CN O

CN N

OR*

OR* 5 “Cation pool”

6

N R

H N O

1, levobupivacaine, R = n-Bu 2, ropivacaine, R = n-Pr 3, mepivacaine, R = Me

1437

Electrosynthesis of Bioactive Materials

Surprisingly, the source of cyanide played a major role in determining the efficiency of the reaction. Initial attempts using NaCN led to complex mixtures and none of the desired product. The use of n-Bu3SnCN afforded a modest amount of 6 (40% yield, 85% ee), while a combination of TMSOTf and TMSCN provided 65% of the desired product but with a diminished 65% ee. Interestingly, the addition of catalytic quantities of β-cyclodextrin saw an increase in both the ee and yield over that recorded for the best nucleophile under otherwise identical reaction conditions. Presumably, the β-cyclodextrin creates a hydrophobic pocket in which the chiral auxiliary, 8-phenylmenthyl, resides (see 7). A subsequent Re-face attack by the nucleophile generated the new stereogenic center with the appropriate absolute configuration required for the synthesis of the natural product. Acid hydrolysis to remove the chiral auxiliary, followed by coupling with dimethyl aniline, and a final N-alkylation afforded the natural product. Thus, the electrolytic cation pool method coupled with biomimetic cocatalysis led to an efficient synthesis of ropivacaine (2).

+

–2e– (100 mA) N CO2R* 4

–

Pt anode/W cathode MeOH, 0°C, 5 h β-cyclodextrin (β-CD) R* = 8-phenylmenthyl

N

HO

HO

O

HO

–

O

N CN CO2R* 6 no β-CD β-CD –

HO

CN

HO

HO

%6 % ee

7

HO

85 76

65 91

Another example of the electrochemical formation and use of acyliminium ions is seen in the example portrayed in the accompanying equation. Pymetrozine (8) is known for its insecticidal properties. Environmentalists interested in its fate and potentially harmful effects of this substance have identified several interesting metabolites including structure 10. Hudlicky and coworkers capitalized upon the electrochemical oxidation of amides and developed an efficient constant current electrolysis (CCE) route to convert 8 to 10 [16]. OMe N N– N–

N H

N

O

8, pymetrozine

+ N –

CCE, Et4NOTs, MeOH (62%)

N N–

N H

O 9

N

N N– N–

N H

N

O 10

2. Tetrahydroisoquinoline Alkaloids Several of the biosynthetic pathways leading to alkaloids involve oxidation at nitrogen, a process that can be achieved electrochemically, photochemically, or by using one-electron oxidizing agents [17]. Hurvois and coworkers used an electrochemical approach to successfully synthesize tetrahydroisoquinoline (THIQ) alkaloids called (±)-carnegine (11), (±)-O,O-dimethylcoclaurine (12), and (±)-norlaudanosine (13) [18]. The key synthetic intermediate, α-aminonitrile 15, was generated via an anodic oxidation of tertiary amine 14 followed by cyanation of the iminium ion 17 that was generated in situ. Voltammetry was used to address whether the catechol or the amine was first to oxidize. A cyclic voltammogram of 14 was obtained using a reticulated vitreous carbon (RVC) anode in a 0.1 M LiClO4/MeOH solution at a sweep rate of 50 mV/s. In the absence of a nucleophile, two irreversible peaks were observed at 0.9 and 1.3 V (vs. SCE). The first oxidation peak is consistent with previous voltammetric studies on a THIQ alkaloid that did not contain a catechol moiety, thereby suggesting that the initially formed radical cation resides on the amine rather than the catechol.

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1438

Organic Electrochemistry

Unfortunately, regioselectivity issues arise once the initial oxidation occurs, because two iminium ions, 17 and 18, are generated from a nonselective deprotonation of 16 [19]. The regiochemical issue was overcome by reducing the number of equivalents of NaCN (from 4–6 to 2–2.5) and by adding 0.5 equiv. of AcOH, the latter presumably serving to scavenge base produced at the cathode during the electrolysis. Alkylation of the lithiated α-aminonitrile derived from 15, followed by reductive decyanation and hydrogenolysis, led to the synthesis of the racemic alkaloids. 0.95 V (SCE), 0.1 M LiClO4, MeOH, 2.5 equiv. NaCN, 0.5 equiv. AcOH

MeO N

MeO

Ph

MeO MeO N

MeO

(85%)

R 11, R = Me 12, R = 4-MeOBn, O,Odimethylcoclaurine 13, R = 3,4-(MeO)2Bn, norlaudanosine

CN

14

15

MeO

+

N

MeO

Ph

+

–H –e

NH

MeO

Ph

MeO

MeO N

MeO

Ph

+

16

N

MeO

+

Ph

+

18

17

III. COUPLINg OF ELECTRON-RICH AROMATICS A.

O-METHYLTHALIBRINE

The diaryl ether called O-methylthalibrine (19) has been used as a folk medicine in central Asia for the treatment of malaria. The retrosynthetic analysis envisioned by Nishiyama and coworkers, and illustrated in the accompanying scheme, visualizes an electro-oxidative dimerization of a dihalophenol, 21, to form diaryl ether 20, followed by installation of the two isoquinolines and cleavage of the carbon-halogen bonds [20]. OCH3 O NMe H

H3CO

OH HO O

Hal

Hal Hal

R Hal

H MeN H3CO

Hal

OCH3

Cleave C-Hal Add isoquinolines

OCH3 19, O-methylthalibrine

R

Electrooxidative coupling

R 21

20

In practice, the dimerization of dichlorophenol 22 was carried out in a glassy carbon beaker that served as the anode, under constant current conditions. Once formed, the quinoid adduct was reduced using a zinc plate cathode to produce phenol 24. It, as well as the corresponding brominated structure, 25, was converted to the advanced stage synthetic intermediate 26. Of the two, only the brominated form, 26b, could be converted to the natural product.

HO

Cl

2 Cl

CCE (10 mA/cm2), glassy carbon beaker, MeOH, 0.1 M HClO4

R 22, R = CH2CO2CH3

© 2016 by Taylor & Francis Group, LLC

(60%)

Cl

O

Cl O

MeO

Cl

R Cl

OH

Zn plate cathode (61–99%)

O R

Cl

R 23

Cl

R 24

1439

Electrosynthesis of Bioactive Materials X

OH

X

O R

X

OMe O

Steps

NMe H

X

X OMe OMe

H2, 5% Pd/C

X

19

(48%) H MeN

R MeO

OMe

24, X = Cl, R = CH2CO2CH3 25, X = Br, R = CH2CO2CH3

26a, X = Cl; 26b, X = Br

IV. ELECTROgENERATED REAgENTS A.

HYPERVALENT IODINE

1. Tetrahydropyrroloiminoquinone alkaloids The tetrahydropyrroloiminoquinone alkaloids have attracted a great deal of attention, largely because of their interesting molecular architecture and cytotoxic activity toward human colon tumor cells, esophageal cancer, and doxorubicin-resistant L1210/DX murine lymphocytic leukemia cells. Nishiyama and coworkers have developed an interesting route to these materials based upon the in situ electrochemical generation of PhI(OCH2CF3)2, a useful hypervalent iodine reagent [21]. The accompanying equations illustrate their approach. Ultimately, the quinolinone adduct 28 was converted to 29 and 30, two members of the tetrahydropyrroloiminoquinone alkaloid class. CCE (10 mA) Phl LiClO4, CF3CH2OH

BnO

NO2 CO2Me CONHOMe

Phl(OCH2CF3)2

NO2 CO2Me

in situ generated Phl(OCH2CF3)2, CF3CH2OH

BnO

(62%)

–

N

27

O

OMe 28 NH

N O HO

N

OH 29, N-1-β-D-ribofuranosyldamirone C

B.

N

O O OH

O HO

NH2 O OH

OH 30, N-1-β-D-ribofuranosylmakaluvamine I

METHOXY RADICAL

1. En Route to Parasitenone Boron-doped diamond (BDD) electrodes have attracted a great deal of attention and have proven to be particularly effective for the generation of such potent oxidizing agents as the hydroxy radical [22]. In early 2012, evidence was obtained for the formation of the methoxy radical from the oxidation of methanol at a BDD anode [23]. Interestingly, the intensity of the ESR signal corresponding to the adduct of 5,5-dimethyl-1-pyrroline-N-oxide (DMPO) and the radical varied as a function of the anode material, with BDD being larger than Pt, and nearly zero for a glassy carbon electrode. Researchers exploited these

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1440

Organic Electrochemistry

findings in a total synthesis of the natural product called parasitenone (31), an inhibitor of interleukin TNF-α. When its precursor, 32, was oxidized at a potential of 1.15 V (SCE) using a platinum net anode, a nearly quantitative yield of the bis-dimethylketal 33 was generated. A similar outcome was obtained when a BDD anode was utilized. In contrast, the use of a glassy carbon beaker type cell/anode provided a 50% yield of aldehyde 34, but none of the bisketal. Also of interest is the stark difference between the electrochemical- and reagent-based protocols, the latter failing to deliver any of the bisketal (Table 37.1). The difference between the electrochemical- and reagent-based outcomes was attributed to the fact that the latter do not generate the methoxy radical. This was also said to be the case when a glassy carbon electrode was utilized. In that case, the chemistry presumably follows the alternative pathway, illustrated in the following scheme, and leading to the formation of aldehyde 34. –e, –H

+

OMe

–e

MeO

MeOH

MeO

+ OMe

OMe OR

–e

–

OR

+

OR

33

MeO OMe 32

–

MeO OMe 36

OMe 35 MeO

MeO

– +

OMe

OMe

OR

OR –e

34 –MeOH

OMe 37

OMe 38

TAbLE 37.1 Reaction of Electrochemically generated Methoxy Radical OMe

MeO

OH

OMe

OTBDPS

OTBDPS [O]

OMe 32

OMe

O O

MeO OMe

31, (±)-parasitenone

33

CHO

OH

OMe 34

Potential (vs. SCE)

Product and Yield (%)

Pt net

1.15

33 (100)

2

Glassy carbon beaker

1.05

34 (50)

3

BDD plate

1.25

33 (100)

4

DDQ

—

34 (100)

5

Phl(OAc)2

—

34 (50)

6

CAN

—

34 (70)

Entry

Anode/Oxidant

1

Notes: 5% KOH MeOH, Pt wire cathode for each electrochemical experiment; DDQ refers to 2,3-dichloro-5,6-dicyanobenzoquinone and CAN to ceric ammonium nitrate.

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1441

Electrosynthesis of Bioactive Materials

V. UMPOLUNg A.

BICYCLO [3.2.1] SKELETON

Both oxidative and reductive routes have been developed to access the bicyclo [3.2.1] framework. Much of the interest in these systems stems from their prevalence in the structure of many natural products. Each route capitalizes upon the unique ability of electrochemistry to achieve umpolung: reduction turns electron-deficient alkenes into nucleophilic intermediates, while oxidation changes an electron-rich alkene into an electrophilic species [24]. Illustrated in Equations 37.1 and 37.2 are examples, the first corresponding to a controlled potential cathodic reductive coupling that was used to complete a formal total synthesis of the antitumor antibiotic called quadrone [25]. The second equation portrays a constant current anodic oxidation reaction that was targeted toward the synthesis of scopadulcic acid B [26], a substance that is said to be useful in the treatment of stomach disorders and improve digestion [27]. CN

OH

–2.4 V (SCE), CH3CN n-Bu4NBr, CH2(CO2Me)2

OTBDPS

OTBDPS

NC

(37.1)

(90%) CHO

RVC anode, Pt cathode 0.1 M LiClO4, 30% MeOH/THF

S

S

OMe 3

S MeO

S

CH(OMe)2

(37.2)

2,6-lutidine, 8 mA, 2.2 F (75%)

B. (−)-CROBARBATIC ACID, (+)-NEMORENSIC ACID As noted previously, oxidation provides a versatile, convenient, and efficient means of achieving umpolung. Many fundamental investigations of both oxidative and reductive processes have been conducted, and the number of applications to the total synthesis of natural products continues to grow. Moeller and coworkers, for example, used anodic oxidation to reverse the polarity of electron-rich alkenes and allow the resulting cation radical to be intercepted by a pendant nucleophile (e.g.,  the conversion of 42 to 41). Application of the concept to the total synthesis of the furancontaining portion of the macropyrrolizidine alkaloid called (+)-nemorensic acid (39) [28], and the related structure called (–)-crobarbatic acid (40) [29], is described in the following text. These substances exhibit hepatotoxicity and antitumor activity [30,31]. Notice that each structure possesses a quaternary center at C-2 and that the orientation of the methyl groups positioned at C-2 and C-3 differs in the two structures. As the ensuing discussion shows, electrochemistry provides an efficient and intriguing approach to solving these challenging problems.

CO2H

3

HO2C

O 2 CO2H

O

O

(+)-nemorensic acid (39) (–)-crobarbatic acid (40)

R3

R1 R2

X΄ OH 41

+

–e

R2 X

R3

R1

X΄ OH

X

42

To begin, the enol ether model systems, 43a–e and 45a,b, were subjected to a CCE (8 mA) in a 30% MeOH/THF solution with 2,6-lutidine as a proton-scavenger using an RVC anode and platinum cathode. As illustrated, five- and six-membered rings formed in yields ranging from 51% to 96% (note structures 44a–e and 46a–e). The cyclization of 45a and b demonstrated that it was possible to generate a quaternary center. Attempts to extend the methodology to the construction of

© 2016 by Taylor & Francis Group, LLC

1442

Organic Electrochemistry MeO

OMe

OMe

OH

n R 43 a–e

44 a–e MeO

OMe

O

n 45 a,b

FIgURE 37.1

R n

0.03 M Et4NOTs, 30% MeOH/THF 2,6-lutidine

OMe

HO

O

CCE (8 mA, 2 F) RVC anode, Pt cathode

a, R = Me, n = 1 (96%, 5:1 trans:cis) b, R = Bn, n = 1 (95%, 10:1 trans:cis) c, R = H, n = 1 (80%, NMR yield) d, R = Bn, n = 2 (51%, 6:1 trans:cis) e, R = H, n = 3 (n/a)

a, n = 1 (74%, 3:1 trans:cis) b, n = 2 (56%, 1:1 trans:cis) 46 a,b

n

Formation of five- and six-membered rings.

seven-membered rings proved unsuccessful; instead, the initially formed cation radical is captured by methanol. Use of a ketene dithioacetal, 47a–d, in place of the enol ether works well, leading to the formation of five- and six-membered rings in high yields and to the efficient formation of quaternary centers. CCE (8 mA, 2.2 F) RVC anode, Pt cathode

R

n HO

R

S

0.1 M Et4NOTs 30% MeOH/THF 2,6-lutidine

S

47 a–d 47

R

n

48 (% yield)

a

H

1

94

b c d

H

2

87

Me

1

83

Me

2

70

R

n

R S O MeO S 48 a–d

With an effective method for generating the stereochemistry and quaternary center needed for the synthesis of (+)-nemorensic acid (39) in hand, a 3:1 mixture of stereoisomeric thioketene acetals 49 was subjected to CCE (8 mA, 2 F) in an undivided cell to obtain a 71% yield of a separable mixture of isomers. The electrolysis product, 50, was subsequently elaborated to afford the natural product in a 65% yield over an additional three steps. S OH

S

49 (3:1 ratio of isomers)

CCE (8 mA, 2 F) RVC anode, Pt cathode 0.03 M Et4NOTs 30% MeOH/THF (71%)

S O MeO S 50

Steps HO2C

O

CO2H

(+)-nemorensic acid (39)

Unlike nemorensic acid (39), the synthesis of crobarbatic acid (40) requires that the vicinal methyl groups be trans to one another. A clever sequence was devised that permitted application of the chemistry just described to the new target structure. It called for an indirect approach wherein the propenyl group found in the cyclized adduct, 52, served as a surrogate for the carboxylic acid found at the quaternary carbon of the natural product. In this case, a postelectrolysis reduction of the thio-orthoester in 52 and oxidative cleavage of the olefin promised to generate a quaternary center with the stereochemistry needed for the assembly of (–)-crobarbatic acid (40) and opposite to that of (+)-nemorensic acid (39).

© 2016 by Taylor & Francis Group, LLC

1443

Electrosynthesis of Bioactive Materials

2

HO

-2e, MeOH

S

O MeO

S

S O S

CO2H O1

(–)-crobarbatic acid (40) 52 alkene = CO2H synthon thio-orthoester = CH3 synthon

51

In principle, the propenyl unit could have stabilized the intermediate cation radical 53 to the point where it was kinetically unreactive. Fortunately, it did not. Thus, a CCE (8 mA, 1.6–1.8 F) of 51 in a 30% MeOH/THF solution with 2,6-lutidine as an acid scavenger afforded 52 in a 57% yield as a 5:1 ratio of diastereomers. The yield improved to 72% when n-butyllithium was used as an acid scavenger for the methanol that was reduced at the cathode. The natural product, (−)-crobarbatic acid (40), was obtained after an additional five steps from the electrolysis product, 52.

51

RVC anode Pt cathode 0.1 M Et4NOTs 30% MeOH/THF n-BuLi (72%, 5:1 mix of diastereomers)

S OH

+ S 53

VI. A.

S O MeO S

Steps (–)-crobarbatic acid (40)

52

TERPENOIDS FURAN-OLEFIN COUPLING

1. Eunicellin Diterpenes Polycyclic systems containing quaternary centers, and more specifically furans, are often useful intermediates in the synthesis of natural products. Wright and coworkers took advantage of electrochemical umpolung, accessing the fused furans, 55, stereospecifically via a facile two-step oxidative annulation (note Table 37.2) [32]. Conjugate addition followed by formation of the TMS enol ether set up a doubly nucleophilic system, 54. Although either olefin could undergo oxidation to generate the electrophilic radical cation, the furan moiety has been employed as a nucleophilic terminator using electrochemical oxidation [33]. Optimal reaction conditions were achieved using CCE of a variety of substrates, using as little as a 0.1 M LiClO4 as the supporting electrolyte and 2,6-lutidine as an acid scavenger. To eliminate competitive methanolysis of the silyl enol ether, 2-propanol was substituted for methanol as the cosolvent with acetonitrile. As illustrated in the table, the most critical aspect of the study was the relationship of current density to product ratio. Interestingly, electrodes that provided a lower current density resulted in a higher product yield with near-theoretical (2 F) charge consumption. An increase in current density resulted in erosion of product formation and considerably larger charge consumption. It is suggested that this is due to the formation of oligomers when higher concentrations of the electroactive intermediate are present at the electrode [34]. Additionally, these reactions provided exclusively cis-fused products for five- and six-membered enones but the stereocontrol decreased for seven-membered rings, presumably due to the flexibility of the larger ring. The utility of the furan as a terminating moiety was demonstrated when the fused ring system, 56, constructed from the two-step oxidative annulation protocol, was subsequently elaborated into precursors for eunicellin diterpenes [35]. 2. Cyathin Core Nerve growth factor (NGF) has sparked interest as a potential therapeutic agent for the treatment of neurodegenerative disorders [36]. Although there has been limited success in the direct

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1444

Organic Electrochemistry

TAbLE 37.2 Role of Current Density O

O

TMSO

O

MgBr Cul, THF TMEDA, TMSCl

n n = 0–2

O MeCN/2-propanol (4:1) 0.4 M LiClO4 0.1 M 2,6-lutidine

n

n

54

55

Electrodes

Current Density (mA/cm2)

Yield (%)

Steel

0.5

68

Carbon

Steel

1.0

66

Carbon

Steel

5.0

35

4

Carbon

Steel

10.0

10

5

Platinum

Carbon

0.1

54

6

Platinum

Carbon

0.5

27

Entry

Anode

Cathode

1

Carbon

2 3

R1 OR

O

CCE (2 F)

O

RO

Diels–Alder

56

R2 O

O O

N N

O OR1 R2 eunicellin diterpenes

administration of NGF [37], small-molecule inducers of NGF production offer a promising alternative [38]. Erinacine C (57) and scabronine A (58) are two such molecules, belonging to the cyathin class of diterpenoid natural products. Wright and coworkers devised a creative electrochemically based approach to the cyathin skeleton, employing an anodic cyclization reaction to generate the AB ring system (note 61) while installing a substituted furan from which the C ring could be formed via a [4 + 3] cycloaddition reaction [39], as illustrated in Scheme 37.1. Initially, silyl enol ether 60 was oxidized using a CCE in an MeOH/MeCN solution. Unfortunately, desilylation rather than cyclization occurred. Use of a less nucleophilic solvent combination wherein isopropanol replaced methanol as the cosolvent with acetonitrile met with success. Indeed, when 60 was oxidized in a CCE (CCE; 100 mA) using a carbon anode in an LiClO 4 solution of 20% i-PrOH/MeCN with 2,6-lutidine added as a proton scavenger, the cis-fused tricyclic product 61 was obtained in 65% yield. 3. Tricholomalide A The diterpene metabolite called tricholomalide A (64) induces the expression of NGF, and may present a potential treatment for neurodegenerative diseases [40]. Wright and coworkers envisioned that the ABC-tricyclic core of the natural product could be accessed through a two-step sequence involving a conjugate addition to the substituted cyclopentenone 67 followed by an anodically initiated annulation to afford 65 [33,41].

© 2016 by Taylor & Francis Group, LLC

1445

Electrosynthesis of Bioactive Materials

Br

CCE (100 mA) carbon cathode

a) Mg, Cul, TMSCl b)

O 59

TBSO

O

2,6-lutidine 0.1 M LiClO4 1:5 i-PrOH/MeCN

O 60

OH O

O

O 61 65% (2 steps) [4 +3] and additional steps

HO2C

R O

O

B

A

O

O

OMe

O

O

MeO

O

OH scabronine A (58)

HO

63, R = CH3, CO2H

62

O

HO erinacine C (57)

SCHEME 37.1

Key intramolecular oxidative coupling to access the cyathin diterpenoids. O O CCE OH O O (64), tricholomalide A

O 65

TMSO

O

66

O

O

67

Model studies proved encouraging. For example, a constant current oxidation of enol ether 69 in a 0.1 M solution of LiClO4 dissolved in i-propanol, using 2,6-lutidine as a proton scavenger, delivered the tricyclic adduct 70 in a respectable 74% over the two-step sequence. The same series of steps, starting with cyclohexenone 71, afforded a 65% yield of the tricyclic framework 73. O

constant current MgBr

O 68

Cul, Et3N TMEDA, TMSCl

TMSO 69

O MgBr

Cul, Et3N TMEDA, TMSCl O 71

TMSO

O 72

O

0.1 M LiClO4 2,6-lutidine, i-PrOH (74% overall) CCE (0.36 mA, ~2 F) 0.1 M LiClO4 2,6-lutidine, i-PrOH (65% overall)

O 70

O

O

O 73

Unfortunately, extension of the methodology to cyclizations leading to the formation of five- and seven-membered rings met with mixed outcomes. Difficulties encountered in the synthesis of the silyl enol ether precursors thwarted efforts to construct five-membered rings, while the successful annulation to form a seven-membered ring depended upon the presence or absence of a substituent at C-3 of the enol ether starting material, 74a,b (note Equation 37.3). Attempts to extend the methodology to thiophene substrates proved satisfactory for the formation of six-membered rings, though

© 2016 by Taylor & Francis Group, LLC

1446

Organic Electrochemistry

seven-membered ring annulations failed to proceed, even in the presence of a quaternary center at C-3, possibly due to the reduced nucleophilicity of the thiophene. R

R

CCE (RVC anode) 0.1 M LiClO4

3

2,6-lutidine i-PrOH/MeCN

TMSO

O 74a (R = H), 74b (R = CH3)

(37.3)

75, R = H (0%) 76, R = CH3 (70%)

O

O

An interesting competition study involving the bis-furanoid structure 77 indicated a clear kinetic preference for formation of six-membered rings. Thus, while cyclization could have afforded either a six- or a seven-membered ring, only the six-membered ring, 78, was produced. Finally, competitive closure between a furan and a thiophene resulted solely in the furyl-terminated product, confirming the reduced nucleophilicity associated with the thiophene. O

O 3

[O] 70%

TMSO

O O 77

78

O

The mechanism of the annulation step was briefly investigated. Voltammetric analyses of model compounds 79 and 80 demonstrated that the silyl enol ether had a lower oxidation potential (E1/2 = 0.87 V vs. Fc/Fc+) than the furan (E1/2 = 1.31 V vs. Fc/Fc+), indicating that the enol ether electrophore was the first to oxidize (Scheme 37.2). It was reasoned that oxidation of the probe structure 81 ought to afford a labile cyclopropyl carbinyl radical cation (85) that should undergo ring opening at a rate that is at least competitive with cyclization. Surprisingly, cyclization dominated and no evidence for opening of the cyclopropane ring was observed. The authors suggest that cleavage of the Si–O bond occurs more rapidly than ring opening to afford an α-keto radical 84 that subsequently closes onto the furyl unit leading to the tetracyclic delocalized radical, 82. An alternative rationale posits that intramolecular electron transfer equilibrium between the initially formed radical cation, 85, and the pendant furan simply reverses the role of the electrophores so that the enol ether serves as the nucleophilic component. 4. (−)-Alliacol A Alliacol A (87) presents an attractive synthetic target. In addition to the challenges presented by the structural complexity of its core (viz., five contiguous stereocenters and three contiguous quaternary carbons), its moderate antitumor and antimicrobial activity add to the appeal [42].

OTMS 79 OTBS

H 1 H O TMSO 81

–e –TMSOR

–e

H H O

O 82

69%

H H O

O 83

O 80 H H O

O 84

SCHEME 37.2

Annulation mechanism.

© 2016 by Taylor & Francis Group, LLC

H + H O TMSO 85

H 1 + H O TMSO 86

1447

Electrosynthesis of Bioactive Materials

Although several racemic syntheses have been reported [43], Moeller and coworkers carried out a direct, asymmetric synthesis using an anodic oxidative cyclization of 89 as the key step to obtain the B-ring of the core [44]. 10

O

O

TBSO

1 Q

O

CCE

B HO

OH

O (–)-alliacol A (87)

TBSO

O

O

88

89

In model studies conducted prior to engaging in the total synthesis effort, the C-1 desmethyl enol ether 90 was examined to establish that the quaternary carbon of the tricyclic core, designated as CQ in structure 87, could be accessed electrochemically. A CCE (12.9 mA, 2.1 F) was employed using an undivided flask using an RVC anode and a carbon rod cathode in a 0.4 M LiClO4 solution of 20% MeOH in CH2Cl2 and with 2,6-lutidine added as a proton scavenger. The initial cyclized product, 91, was immediately treated with p-toluenesulfonic acid to generate the furan, 92, in yields ranging from 85% to 90%. O

OTBS

O

CCE (12.9 mA, 2.1 F) TBSO

RVC anode, carbon cathode 2,6-lutidine, 0.4 M LiClO4 1:5 MeOH/CH2Cl2

O

B TBSO

O

p-TsOH B (85–90%)

RO

MeO 91

90

O

92 R = H or TBS

Once it was determined that the bicyclic skeleton could be generated electrochemically, both chemical and electrochemical protocols were explored to generate the A-ring, and establish the quaternary center (Scheme 37.3). While radical cyclization methods (viz., AIBN with n-Bu3SnH or (n-Bu3Sn)2) proved unsuccessful, anodic oxidation of the enol ether 93 derived from 92 did achieve cyclization to produce 94 in a respectable 74% yield. Despite the success of this transformation, the presence of a superfluous formyl group equivalent (R=CH(OMe)2) on the A-ring of 94 detracted from the sequence and led to the use a Friedel–Crafts alkylation reaction in its place (95–96). O

O CCE (19.7 mA, 2.1 F) (74%)

O MeO

MeO 94, R = CH(OMe)2 O

93

94

O

I2, PPh3 imidazole 80%

A RO

AgNO3, MeOH I

THF, 71%

O

MeO 96

95 TBSO TBSO

O 97

SCHEME 37.3

CH3 a) CCE (15.3 mA, 2.1 F) b) p-TsOH HO

O

O

CH3 a) I , PPh , 2 3 imidazole b) AgNO3 MeOH/THF

98

Studies directed toward the synthesis of alliacol A.

© 2016 by Taylor & Francis Group, LLC

A O

CH3

O O MeO

99

1448

Organic Electrochemistry

With a solid foundation in hand, the anodic cyclization of enol ether 97, a system differing from the model by inclusion of the pro-C-1 methyl group, was examined for the construction of 98 and its conversion to the tricyclic framework 99. Ultimately, both a racemic and an asymmetric synthesis of alliacol A (87) were accomplished using a tandem anodic oxidation and Friedel–Crafts cyclization to arrive at the natural product. 5. Arteannuin Skeleton The arteannuin family of natural products is isolated from the same leaves that produce artemisinin, a novel and potent antimalarial drug. The arteannuins are believed to be key intermediates in the biosynthesis of artemisinin [45]. In 1996, Schwaebe and Little examined the cathodic reduction of keto enoate 102, in hopes that the product of electrohydrocyclization could be transformed to the desmethyl analogue of arteannuin B (100) [46,47]. While cyclization occurred in a respectable yield (78–95% of 103 from 102), the process was not stereoselective. In contrast, treatment of 104 with samarium diiodide in THF/ MeOH at 0°C afforded the γ-hydroxy ester 105 with the stereochemistry needed to convert it to the natural product analogue 100. In this instance, samarium was believed to exert a templating effect by complexing with Lewis basic sites to position the reacting centers as illustrated by structure 106. H

H CO2Me

3

O

O

3

HO

102 H

HO CH2CO2Me 103

CO2Me

O

HO CH2CO2Me 105

104

R

‡

1

100, R = H (desmethyl arteannuin B) 101, R = CH3 (arteannuin B)

O O

XSm

O

O O OMe 106

Capitalizing on the success of the electrochemical umpolung reaction in the synthesis of alliacol A (87), Moeller and Wu examined an anodic oxidation pathway to the synthesis of the arteannuin skeletal framework exemplified by structures 109 and 110 [48]. The approach is outlined in Scheme 37.4. At the outset, it was not known whether the methyl group at the four-position on the C-ring (ultimately destined to become the α-methylene unit in the natural product) would affect the reactivity of the furan; thus, the synthesis was initially carried out both in the absence, 111a, and presence, 111b, of a methyl group on the furan. In the former case, the methyl group would be installed at a later stage in the synthesis. The first anodic oxidative cyclization reaction was performed using the same reaction conditions described in the synthesis of alliacol A. In an undivided three-neck flask, 111a/b was oxidized using an RVC anode and a carbon rod cathode under CCE conditions (20 mA) until 2.1 F of charge had passed. In this synthesis, however, the supporting electrolyte was changed to a 0.1 M Et4NOTs solution in 20% MeOH/CH2Cl2 with 2,6-lutidine added as a proton scavenger.

Anodic oxidation

X O

OH

O alliacol A (87)

SCHEME 37.4

O

H A

X O

R R 107 R = H and R = CH3 108

Intramolecular oxidative coupling.

© 2016 by Taylor & Francis Group, LLC

10

O

R’3SiO

O

O OC

H

1

B H

O 109, arteannuin B

HO O H HO O 110, arteannuin M

1449

Electrosynthesis of Bioactive Materials

Interestingly, the unmethylated furan, 111a, produced 60% of the desired cyclization product, 112a, while the methylated furan, 111b, generated a meager 30% of 112b. It was postulated that 4-methyl substituted furan 111b (Ep/2 ~ +1.46 V vs. Ag/AgCl) was oxidizing competitively with the enol ether initiating group (Ep/2 ~ +1.44 V vs. Ag/AgCl) [49] as compared to 111a (Ep/2 ~ +1.64 V vs. Ag/AgCl), thereby accounting for the lower yields associated with the methyl case. To circumvent this issue, an N,O-ketene acetal initiating group (Ep/2 = +1.15 V vs. Ag/AgCl) was used in place of the enol ether. The modified substrates, 113a,b, were subjected to the same oxidation conditions and produced the desired product, 114a,b, in excellent yields. OMe TIPSO

TIPSO

CCE (20 mA, 2.1 F ) O

4 R 111a, R = H 111b, R = Me

O

RVC anode, carbon cathode 2,6-lutidine, 0.1 M Et4NOTs 1:5 MeOH/CH2Cl2

O TIPSO N O O Ph

4 R 112a, R = H (60%) 112b, R = Me (30%) O

O Same conditions

O

N O Ph

4 R 114a, R = H (70%, 87% by NMR) 114b, R = Me (65%, 90% by NMR)

4 R 113a, R = H 113b, R = Me

The bicyclic adduct 114a,b was subsequently converted into 115a,b in preparation for a second electrochemically promoted cyclization. Similar electrolysis conditions to the first anodic oxidation were employed (8.0 mA, 2.4 F) with the exception that LiClO4 was used as the supporting electrolyte for ease of removal during the aqueous workup. The tricyclic core of the arteannuin skeleton, 116, was obtained in excellent yield independent of the presence of the methyl group on the furan ring, thus demonstrating the versatility of the electrochemical cyclization. O O

O N PhO

R 114a, R = H 114b, R = Me

H

Steps

MeO

CCE (8 mA, 2.4 F) O O RVC anode, carbon cathode 2,6-lutidine, 0.1 M LiClO4 R R O 1:5 MeOH/CH2Cl2 115a, R = H 116a, 70% (87% by NMR) 115b, R = Me 116b, 70% (80% by NMR) O

6. (−)-Heptemerone b and (−)-guanacastepene E Heptemerone B (117) and guanacastepene E (118) possess a unique tricyclic skeleton that is characteristic of the neodolastane class of diterpenes. Unlike the guanacastepenes that display activity against antibiotic-resistant bacteria [50], the heptemerones demonstrate little antibiotic activity. Rather, they are potent inhibitors of fungal germination for Magnoporthe grisea, the cause of rice blast disease [51]. The related structures and unique biological activity of heptemerone B (117) and guanacastepene E (118) have made them interesting synthetic targets. In 2006, Trauner and coworkers synthesized these compounds using a convergent approach [52]. As the ensuing discussion illustrates, they elegantly demonstrated the usefulness of anodic cyclization to generate a complex, late-stage intermediate that was used to complete the total synthesis of the natural products. The starting material for the oxidation viz. furyl ketone 119 was produced via a cuprate coupling reaction between the fragments, 121 and 122, of the core framework; 119 was then converted to the silyl enol ether, 120,

© 2016 by Taylor & Francis Group, LLC

1450

Organic Electrochemistry

PO

–e 2

PO

1

Cyclization

O

OBn

PO

+

O OTBS 120 (P = TBDPS)

O H +

TBSO BnO

OBn OTBS

124

123

CCE (0.9 mA), 0.1 M LiClO4, 20% MeOH in CH2Cl2 (81%) PO

PO MeO

O H

OBn MeO

O

125

126

SCHEME 37.5

O H

OBn OTBS +

Key electrooxidative coupling en route to the guanacastepenes.

using a regioselective deprotonation and trapping of the enolate with TBSOTf. Initial attempts to accomplish the annulation using nonelectrochemical methods/techniques were unsuccessful [53]. It was reasoned that bond formation between C-1 and C-2 in 120 could be achieved by oxidation to form a radical cation. Although several reagent-based methods could have been used [54], Trauner employed the electrochemical methods developed by Moeller [55] and Wright [56] to couple the silyl enol ether and furan subunits of 120.

2

PO

OBn

1

O

O

I

119

RO O H

OBn O

O

117, R = Ac, (–)-heptemerone B PO 118, R = H, (–)-guanacastepene E

2 1

O

OBn OTBS

OBn PO 121

O 122 (P = TBDPS)

120 (P = TBDPS)

The anodic cyclization leading to the desired 7-membered ring, 126, was accomplished in good yields (81%) by employing CCE conditions (0.9 mA, 2.6 F) using an RVC anode and platinum cathode, 2,6-lutidine as a proton scavenger, and a 0.1 M LiClO4 solution in 4:1 CH2Cl2 (DCM):MeOH. Trauner proposed that the anodic cyclization began with oxidation of the silyl enol ether to generate 123 (Scheme 37.5). Next, a stereoselective attack of the furan moiety led to the cyclized intermediate, 124. Trapping of the oxonium ion by methanol, oxidation, and subsequent desilylation produced the product, 126. Ultimately 126 was used to complete the total synthesis of (–)-heptemerone B (117) as well as a total synthesis of (–)-guanacastepene E (118).

VII. A.

MEDIATORS [57] DAUCENE

Oxidation of tris(4-bromophenyl)amine (M) at a potential of ca. 0.9 V vs. Ag/0.1 M AgNO3 leads to its well-known and frequently used triarylaminium cation radical [58]. This system has been shown to be capable of mediating the oxidation of a variety of strained hydrocarbons including those of the bicyclo [2.1.0] framework exemplified by structure 127 [59]. The resulting cation radical 128 then undergoes a Wagner–Meerwein rearrangement involving migration of the endocyclic carbon,

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1451

Electrosynthesis of Bioactive Materials

TAbLE 37.3 Rearrangement of Strained Hydrocarbons O

O

CH3

O (p-BrC6H4)3N (20 mol% x 2)

O

Conditions 132 Entry

CH3

CH3

CH3

131, daucene

133 Oxidant

Solvent

Base

Yield (%)

1

(p-BrC6H4)3N . SbCl6

CH2Cl2

2,6-di-t-butylpyridine (0.5 equiv.)

85

2

+0.88 V, 0.1 M TBABr (Pt anode/RVC cathode)

CH3CN

none

70

Cen, toward the center that is best able to stabilize a positive charge. A one-electron transfer from the neutral form of the mediator to the hydrocarbon cation radical formed after rearrangement serves to regenerate the oxidant and produce the product, 130. The mechanism just described and portrayed in the following text predicts that the chemistry ought to occur without the consumption of charge since following the initial oxidation at the anode, the redox mediator is regenerated catalytically. In practice, this is generally not the case. The inefficiency can be traced to several probable causes including reduction of the initially formed aminium cation radical at the cathode in competition with electron transfer between it and the substrate. Ren

Rex

Ren

R΄

Rex R΄

+

R M M 127 M = (4-BrC6H3)3N

R 128

+

Ren

+

[Ren]~ R 129

Ren

R΄

R΄

Rex

Rex M

M + 130

R

This chemistry has been applied to the total synthesis of a sesquiterpene called daucene (131) [60]. The featured step was the rearrangement of the spirocycle 132 to the bicyclo [5.3.0] framework found in the natural product. Unfortunately, a total of 40 mol% of tris(4-bromophenyl)amine, added in two successive batches of 20 mol%, was required in order to completely consume the starting material. Generally, only 10 mol% of the mediator is sufficient to achieve the conversion of the [2.1.0] framework to the rearranged adduct. The rearrangement step was also examined using the commercially available reagent tris(4-bromophenyl)aminium antimony hexachloride in the presence of 2,6-di-t-butylpyridine. As Table 37.3 illustrates, the two protocols delivered the desired product, 133, in comparable yields.

VIII.

OXIDATIVE DEAROMATIZATION

A. INTRODUCTORY REMARKS Oxidative dearomatization is characterized by a multistep sequence whose first phase involves the oxidation of an electron-rich aromatic, 134/137 [61]. Subsequent loss of a proton and a second electron leads to 135/138. The reaction of 135 with a nucleophile leads to the substitution product 136,

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1452

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while a cycloaddition reaction involving 138 leads to 140. Thus, oxidative dearomatization reactions can be recognized by either of these signature reactions. R

R a) –e b) –H

+

Nu

NuH c) –e

OH 134

RO

a) –e b) –H

RO

OH c) –e

R –H+

+

O 135

NuH

O 136 RO

RO

RO +

+

RO

O

+

137

O

O

RO

138

O

139

140

The process is generally accomplished through the use of a hypervalent iodine reagent such as phenyliodine(III) bis(trifluoroacetate) (PIFA), or electrochemically. Examples of each are portrayed in the accompanying equations. The first illustrates a transformation that was investigated in an effort to elucidate the role of redox chemistry in the expression of bioactivity for the pseudopterosin class of marine natural products [62]. The second equation portrays a key step in the total synthesis of aeroplysinin (144) and aerothionin (145), an interesting spiroisoxazoline dimer that displays antimicrobial and cytotoxic activities [63]. The third reaction illustrates the utility of cycloaddition chemistry for the rapid construction of complex ring systems possessing the tricyclic framework of structure 150 [64]. H

H Phl(OCOCF3)2, 0 °C O CH2NO2, CH2CI2 (56%)

O OH

H3C

OH 141, R = CMe2 O 142, R = H, iso-PsE

O

OR OR MeO2C

N O

HON OH

CCE n-Bu4NCIO4, MeCN

Br

(68%)

Br OMe

146

147 O N

OH OH

O Br

Br OMe 144, (–)-aeroplysinin 1

O N H

HO O

R

N H

N O OH Br

Br MeO

OH BnO 148

Br

Br OMe 145, R = CH2CH2 aerothionin

H O

OMe

CCE (10.8 mA) ~2F n-BuNBF4, Ac2O (68%) AcO

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OR

CN

Br

OMe

MeO

OR

O

O 143, R = CMe2

CO2Me

Br

H 3C

HO2C

BnO MeO O 149 (S):(R) = 2:5

Steps OAc

CO2Me 150, 2-epi-cedrene isoprenologue

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Electrosynthesis of Bioactive Materials

B.

HELIANNUOL E

The sesquiterpene called heliannuol E (151) is one of many unique allelopathic agents isolated from sunflowers. Utilization of its allelopathic activity could provide an alternative to the environmentally toxic herbicides currently used [65]. Nishiyama and coworkers completed both a racemic [66] as well as an enantioselective synthesis [67] of the natural product. Their approach capitalized upon their earlier findings [68] wherein electrochemically generated spirodienones, 153, undergo ring expansion to form dihydrobenzopyrans, 154, when subjected to Lewis acidic conditions. Br

Br

O

HO

OAc

OH AcO

O

Anodic oxidation

OH

Br

152

OAc

O

O

153

HO

HO

OH

OH

OH

heliannuol E (151)

154

Exploratory investigations focused upon the CCE of simple model systems 155 and 156. The studies utilized a glassy carbon beaker as the working anode, a platinum wire as the cathode, and LiClO4 (0.1 M) as the supporting electrolyte; the solvent was varied between methanol and dioxane, both in the presence and absence of aqueous perchloric acid. When X=H, cyclization was inefficient and delivered only low yields of product 157. When a second bromine was appended to the aromatic ring, however, cyclization occurred to afford the spirocyclic product, 158, in a 50% yield. X

X HO

OH

CCE (1.6–1.8 V vs. SCE, 8.2 F) GC beaker anode, Pt cathode

Br

5:1 dioxane/60% aq. HClO4 0.1 M LiClO4

155, X = H 156, X = Br

O O Br

(37.4)

157, X = H (”low yield’’) 158, X = Br (50%)

The substance that was ultimately used in the total synthesis was produced in relatively low yields when the conditions illustrated in Equation 37.4 were applied to 152, possibly due to side reactions brought on by the presence of the acid-labile tertiary hydroxyl group present on the alkyl side chain. Therefore, several less acidic media were examined; acetone ultimately worked best, leading to the desired product, 153, in a 61% yield. In this solvent, only 2 F were required to achieve complete conversion to the product as compared to the 12 F required when dioxane/60% aqueous HClO4 was employed. Subsequent treatment of 153 with BF3∙Et2O resulted in the rearranged product as a mixture of regioisomers 154 and 159. Chroman isomer 154 was elaborated into (±)-heliannuol E (151) in two additional steps. The enantioselective synthesis was accomplished in a similar manner, starting with 152*. CCE (1.3–1.5 V vs. SCE, 2 F) GC beaker anode, Pt cathode 152

153 acetone, 0.1 M LiClO4 61%

Br

BF3 ·Et2O 75% OAc

OAc

HO

HO +

151 Steps

O 154

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OH

Br

O 159

OH

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Organic Electrochemistry Br

Br

HO OH AcO

CCE (1.3–1.6 V, 2 F) 0.1 M n-Bu4NClO4, acetone

OH

O

Br

OAc BF3 Et2O

O

(68%)

O

(61%)

152*

153*

OAc

HO

OH

OH

154*

IX. CyCLOADDITION REACTIONS A.

QUINONE METHIDES

1. Euglobals Tada and coworkers describe a novel synthesis of six euglobals, 161–166, using an electrochemical oxidation mediated by 2,3-dichloro-5,6-dicyanobenzoquinone (DDQ) [69]. These naturally occurring terpenoids inhibit activation of the Epstein–Barr virus [70] and are believed to be biosynthesized from a common intermediate, grandinol (160), through a hetero-Diels–Alder reaction between an o-quinone methide and α-pinene, β-pinene, or α-phellandrene. R

R HO

HO

O

HO

R΄

R΄

H OH euglobal-G1 (161), R = COCH2CH(CH3)2, R΄ = CHO euglobal-G2 (162), R = CHO, R΄ = COCH2CH(CH3)2

R O

O

R΄

OH euglobal-G3 (163), R = COCH2CH(CH3)2, R΄= CHO euglobal-G4 (164), R = CHO, R΄ = COCH2CH(CH3)2

H OH euglobal-T1 (165), R = COCH2CH(CH3)2, R΄ = CHO euglobal-Ilc (166), R = CHO, R΄ = COCH2CH(CH3)2

Initially excess DDQ was used to generate the o-quinone methide, but the substrate scope was limited. This route provided a variety of decomposition products in addition to undesired cycloaddition products. However, the use of a [poly-(tetrafluoroethylene)]-fiber (PTFE)-coated electrode created a hydrophobic environment in which the polar phenolic substrates were selectively oxidized near the electrode, leaving the more nonpolar terpene substrates positioned close to the fibers. When grandinol 160 was oxidized at 0.45 V vs. SCE using a Pt cathode and fiber-coated glassy carbon anode, a mixture of euglobals 165 and 166 was obtained in a combined yield of 79%. The authors suggest that electron transfer at the anode during the DDQ/DDQH2 redox cycle from the phenolic substrate is responsible for in situ generation of the o-quinone methide, as no product is observed in the absence of the mediator. Subsequent [4 + 2] cycloaddition with the corresponding terpene that remained on the outside of the fibers afforded the natural products. CHO HO

O

OH

OH

CCE (0.45 V, 2.5 F) Et4NOTs, CH3NO2

HO

DDQ DDQH2 (0.2 equiv.)

O

160, grandinol

CHO

euglobal-T1 (165; 51%) + euglobal-Ilc (166; 28%)

O

OH

167 PTFE fiber

B. AzAQUINONES 1. Neuroprotective Agents Just as quinone methides can be generated electrochemically, so too can azaquinones, 170. Their subsequent inverse electron Diels–Alder reaction with enamines provides an elegant and very useful

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Electrosynthesis of Bioactive Materials

synthetic route to polyfunctionalized 1,4-benzoxazine derivatives [71], many of which display neuroprotective properties [72]. The process is carried out via the controlled potential oxidation of aminophenol 168 in the presence of an excess of the amine, R1R2CHCH2NH2. As the accompanying equation illustrates the azaquinone formed in the initial oxidation then serves as the diene toward the in situ generated enamine 169. Of the many systems that were synthesized and screened, structure 172 displayed sufficient bioactivity to be considered a potential candidate for the treatment and prevention of cerebral palsy. H

O

O

H2N

R1

CPE, MeOH, Et4NClO4 excess R1R2CHCH2NHR3

Ph

R2

168

COPh

O

O O

170 OH

H R2 N 1 R

O 172, R =

O Ph

169 OH

H N

H

HN

NHR3

HO

Ph Ph Me N R

O

R3HN

COPh

O 171

X. CARbOHyDRATE ANALOgUES AND PRECURSORS A.

REDUCTIVE DEHALOGENATION: INOSITOL AND CONDURITOL SYNTHONS

The conduritol and inositol carbohydrate analogs are potential inhibitors of glycosidic enzymes, thus making them of interest as antiviral agents [73]. Previous syntheses of these materials have employed metal-based methods (e.g., n-Bu3SnH/AIBN) to reduce vinyl halide precursors such as 173 to afford 174 [74]. Hudlicky and coworkers successfully developed a green, electrochemical alternative that proved superior to the metal-based methodology in several instances (Table 37.4) [75].

TAbLE 37.4 Cleavage of Carbon–Halogen bonds Entry

Substrate

Applied Potential (V) Yield (%)

Entry

Substrate

O

–3.2

62

4

–2.2

73

–2.4

54

OH 178 I OH

O O

HO

–3.0

0

5

OH Br 179

OH 176 Br

Br O

3

50

O

HO

OH 175 Cl 2

–3.0

O

O HO

Yield (%)

I

Br 1

Applied Potential (V)

O

HO OH 177

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–3.0

57

6

O

O

Ph

O

Ph

O 180

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Organic Electrochemistry

The most problematic aspect of the research was said to be the difficulty encountered in separating the products from the supporting electrolyte [76]. X

HO

O

n-Bu3SnH/AlBN

O

or CPE, Hg pool cathode Pt foil anode, Ag/Ag+ reference

OH 173, X = halogen

O O

HO OH

174, protected conduritol F

Precursors to conduritols C, E, and F were synthesized and subjected to controlled potential electrolysis (CPE) in a divided cell with a Hg pool working electrode, Pt foil counter electrode, and an Ag/Ag+ reference electrode in a 0.1 M Et4NBr or n-Bu4NBF4 solution in acetonitrile (Table 37.4). The products of the electrochemical reduction were isolated in slightly lower yields (50–62%) than when a tin hydride radical–initiated reduction was employed (67–85%). The chlorinated substrate, 176 (entry 2), failed to produce any dehalogenated product from the electrolysis, while the use of n-Bu3SnH generated the desired product in a modest 36% yield. On the other hand, the electrochemical protocol proved superior for differentiating between the C-I and C-Br bonds in structure 179 (entry 5), thereby permitting the selective cleavage of the C-I bond (73% electrochemically; 52% using tin hydride). Also of interest is the fact that reduction of the cinnamyl-protected framework 180 (CPE, –3.2 V) using a gold foil working electrode removed the protecting groups and cleaved the carbon-bromine bond, thus providing a mild diol or alcohol deprotection strategy. In contrast, a tin hydride–based reduction led to intramolecular radical cyclization onto the pendant alkene of the cinnamyl unit. This electrochemical methodology was later applied to the synthesis of conduritol F and muco-inositol oligomers, both of which displayed inhibition against several glycosidases [77]. Finally, we note that vinyl bromides were electrochemically reduced in the presence of epoxide and aziridine rings which, in conjunction with a ring-opening step, led to key precursors of (–)-pinitol [78] and amino-inositol synthons [79]. CPE (–2.4 V) of 181 and 182 in an n-Bu4NOH and MeOH solution generated the desired ring-opened products 183 and 184 in 54% yield and 98% yield, respectively [75]. Br

X

O

CPE (–2.4 V vs. Ag/Ag+)

O

Hg pool cathode, Pt anode n-Bu4NOH/MeOH

181, X = O 182, X = NTs

B.

O O

MeO XH

183, X = O (54%) 184, X = NTs (98%)

GLYCALS

The enol ether form of a pyranoside, otherwise known of as a glycal, is easily constructed electrochemically. These substances frequently serve as building blocks in the construction of natural products, C-glycosides, and oligosaccharides. In 2001, Parrish and Little reported that a simple CCE carried out in an undivided cell, using an RVC cathode and a consumable zinc anode, provided a convenient means of converting glycosyl bromides to glycals [80]. The outcome of three of the ten cases that were presented in the original manuscript is highlighted in the following text. Once formed, the glycals can be transformed into 1,2-anhydro sugars by treatment with dimethyl dioxirane. Their subsequent reaction with titanocene dichloride and manganese, in the presence of an electron-deficient alkene, leads smoothly to α-C-glycosides [81]. Here, cyclic voltammetry once

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Electrosynthesis of Bioactive Materials

again displayed its utility, this time as a screening tool to predetermine which alkenes would serve as a useful trap for the anomeric radical that is produced in situ. RVC cathode, consumable Zn anode, 0.5 M LiClO4, THF

O AcO

Br

Undivided cell, CCE (44–90%)

O

AcO AcO

O

AcO AcO

AcO O

O O

OAc (68%)

(47%)

(90%)

C.

AcO

OBn

AcO AcO AcO

O

OAc

VALIENAMINE ANALOGUES

The valienamine core of acarbose (185) is believed to be responsible for its biological activity as an α-glucosidase inhibitor used for the treatment of type II diabetes [82]. Tillequin and coworkers sought to synthesize analogs in an effort to explore new α-glucosidase inhibitors [83]. In attempt to obtain the desired bicyclic core with the correct amino and diol substitution pattern, the oxa ring of 187 was opened using a Ritter reaction [84]. However, the traditionally acidic conditions required for this transformation led to the dehydration products 188 or 189 [85]. HO

HO

HO HO

HO NH2

HO

OH 186, valienamine

H N OH HO

O

OH

HO O HO

acarbose (185)

NO2

O HO

OH

NO2 CF3SO3H

O

CH3CN NHAc

O2 N

O2N 188

OH

HO O HO

NO2 (CF3SO2)2O

O2N

O

187

189

Therefore, a new approach to the synthesis of these bicyclic analogs using an electrochemically promoted Ritter reaction was explored. It proved superior to the chemical approach in part, because the electrogenerated acids (EGAs) formed at the anode were effective at creating a net neutral reaction medium while allowing the acid-promoted reaction to take place. Thus, when 187 was oxidized at a platinum anode using CPE (+2.5 V) until 2.5 F of charge had passed, the desired ring-opened product, 193, was isolated in an 8:2 cis/trans ratio. The elimination products isolated from the chemical approach were not observed using the EGA method. It was postulated that reaction proceeds through the carbonium intermediate, 190, which is in equilibrium with 191, and that the driving force of the reaction is likely the formation of the cis-fused oxazoline ring in 192. In summary, the installation of an amino functionality to form valienamine skeletons in a stereocontrolled fashion was accomplished by an electrochemically modified Ritter reaction.

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Organic Electrochemistry

This methodology highlighted the value of electrochemistry as a viable alternative in organic synthesis, as this transformation was seemingly inaccessible using otherwise harshly acidic, chemical conditions.

187

MeCN, Pt anode (56%)

NO2

NO2

NO2

CPE (2.7 V, 2 F) 0.05 M LiClO4

NO2

+ +

O2N OH 190

O2N

N OH 191

O2N

7 +

N O

H 192

O2N 6 NHAc OH 193

XI. CLOSINg REMARKS We hope that the reader has enjoyed reading of some of the advances that have appeared in the arena of electro-organic synthesis as it was applied to the development of new synthetic methods and to the synthesis of natural products between the years of 2001 and 2011. Previously, we wrote [10]: “It is hoped that some day the methodology will be embraced by all practicing organic chemists, and that it will be properly recognized as a routine tool that is able to assist them in solving problems of both the mechanistic and synthetic varieties.” We are delighted to note that significant progress has been made toward that end. We look forward to an even brighter future.

ACKNOWLEDgMENTS We express gratitude to Amgen and Clorox for supporting some of our research into the area of organic electrochemistry. RG is grateful to the PIRE-ECCI Program, administered by the US National Science Foundation for a fellowship.

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56. (a) Whitehead, C. R.; Sessions, E. H. Ghiviriga, I.; Wright, D. L. Org. Lett. 2002, 4, 3763–3765; (b) Wright, D. L.; Whitehead, C. R.; Sessions, E. H. Ghiviriga, I.; Frey, D. A. Org. Lett. 1999, 1, 1535–1538; (c) Sperry, J. B.; Whitehead, C. R.; Sessions, E. H.; Ghiviriga, I.; Wright, D. L. J. Org. Chem. 2004, 69, 3726–3734. 57. Francke, R.; Little, R. D. Chem. Soc. Rev. 2014, 43, 2492–2521. 58. Wu, X.; Davis, A.; Lambert, P. C.; Steffen, L. K.; Toy, O.; Fry, A. J. Tetrahedron, 2009, 65, 2408–2414. 59. Gerken, J. B.; Wang, S. C.; Preciado, A. B.; Park, Y. S.; Nishiguchi, G.; Tantillo, D. J.; Little, R. D. J. Org. Chem. 2005, 70, 4598–4608. 60. Park, Y. S.; Little, R. D. J. Org. Chem. 2008, 73, 6807–6815. 61. Green, J.; Pettus, T. R. R. J. Am. Chem. Soc. 2011, 133, 1603–1608. 62. Zhong, W.; Little, R. D. Tetrahedron Sympos. Print 2009, 65, 10784–10790. 63. (a) Nishiyama, S.; Yamamura, S. Tetrahedron Lett. 1983, 24, 3351–3352. (b) Ogamino, T.; Nishiyama, S. Tetrahedron 2003, 59, 9419–9423. (c) Ogamino, T.; Obata, R.; Nishiyama, S. Tetrahedron Lett. 2006, 47, 727–731. 64. (a) Takakura, H.; Toyoda, K.; Yamamura, S. Tetrahedron Lett. 1996, 37, 4043–4046. (b) Chiba, K.; Fukuda, M.; Kim, S.; Kitano, Y.; Tada, M. J. Org. Chem. 1999, 64, 7654–7656. (c) Kim, S.; Kitano, Y.; Tada, M.; Chiba, K. Tetrahedron Lett. 2000, 41, 7079–7083. 65. Macias, F.; Varela, R.; Torres, A.; Molinillo, J. Tetrahedron Lett. 1999, 40, 4725–4728. 66. Doi, F.; Ogamino, T.; Sugai, T.; Nishiyama, S. Synlett. 2003, 3, 411–413. 67. Doi, F.; Ogamino, T.; Sugai, T.; Nishiyama, S. Tetrahedron Lett. 2003, 44, 4877–4880. 68. Mori, K.; Yamamura, S, Nishiyama, S. Tetrahedron Lett. 2001, 57, 5533–5542. 69. Chiba, K.; Arakawa, T.; Tada, M. J. Chem. Soc. Perkin Trans. 1, 1998, 2939–2942. 70. Takahashi, M.; Konoshima, T.; Fujita, K.; Yoshida, S.; Nishimura, H.; Hokuda, H.; Nishino, H.; Iwashima, A.; Kozuka, M. Chem. Pharm. Bull. 1990, 38, 2737–2739. 71. (a) Largeron, M.; Fleury, M-B. J. Org. Chem. 2000, 65, 8874–8881. (b) Largeron, M.; Neudorffer, A.; Vuilhorgne, M.; Blattes, E.; Fleury, M.-B. Angew. Chem. Int. Ed. 2002, 41, 824–827. (c) Largeron, M.; Neudorffer, A.; Fleury, M. -B. Angew. Chem. Int. Ed. 2003, 42, 1026–1029. 72. Blattes, E.; Fleury, M.-B.; Largeron, M. Electrochim. Acta 2005, 50, 4902–4910. 73. (a) Zitzmann, N.; Mehta, A.; Carrouee, S.; Butters, T.; Platt, F.; McCauley, J.; Blumberg, B.; Dwek, R.; Block, T. Proc. Natl. Acad. Sci. USA 1999, 96, 11878–11882; (b) Lillelund, V.; Jenson, H.; Liang, X.; Bols, M. Chem. Rev. 2002, 102, 515–554. 74. Hudlicky, T.; Luna, H.; Olivo, H.; Andersen, C.; Nugent, T.; Price, J. J. Chem. Soc. Perkin Trans. 1 1991, 2907–2917. 75. Hudlicky, T.; Claeboe, C.; Brammer, L.; Koroniak, L; Butora, G.; Ghiviriga, I. J. Org. Chem. 1999, 64, 4909–4913. 76. Yoo, S. J.; Li, L.-J.; Zeng, C-C.; Little, R. D. Angewandte Chemie, Int. Ed. 2015, 54, 3744–3747. 77. Freeman, S.; Hudlicky, T. Bioorg. Med. Chem. Lett. 2004, 14, 1209–1212. 78. Hudlicky, T.; Price, J.; Olivo, H. Synlett. 1991, 645–646. 79. (a) Desjardins, M.; Lallemand, M.; Freeman, S.; Hudlicky, T. J. Chem. Soc. Perkin Trans. 1 1999, 621– 628; (b) Oppong, K. A.; Hudlicky, T.; Yan, F.; York, C.; Nguyen, B.V. Tetrahedron 1999, 55, 2875–2880. 80. Parrish, J. D.; Little, R. D. Tetrahedron Lett. 2001, 42, 7371–7374. 81. Parrish, J. D.; Little, R. D. Org. Lett. 2002, 4, 1439–1442. 82. Kameda, Y.; Asano, N.; Yoshikawa, M.; Takeuchi, M.; Yamaguchi, T.; Matsui, K.; Horii, S.; Fukase, H. J. Antiobiot. 1984, 37, 1301–1307. 83. Le Goanvic, D.; Lallemand, M.; Tillequin, F.; Martens, T. Tetrahedron Lett. 2001, 42, 5175–5177. 84. Hammerich, O.; Parker, V. D. J. Chem. Soc. Chem. Commun. 1974, 275–276. 85. Sader-Bakaouni, L.; Charton, O.; Kunesch, N.; Tillequin, F. Tetrahedron 1998, 54, 1773–1782.

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38

Stereochemistry of Organic Electrode Processes Toshio Fuchigami and Shinsuke Inagi

CONTENTS I. II.

Introduction ......................................................................................................................... 1463 Stereocontrol in Electrochemical Reactions ....................................................................... 1463 A. Steric and Polar Factors ............................................................................................... 1463 B. Conformational and Configurational Stabilities.......................................................... 1464 C. Thermodynamic and Kinetic Controls ........................................................................ 1464 D. Stereoselective Synthesis ............................................................................................. 1465 III. Stereoselective and Stereospecific Electroreduction ........................................................... 1465 A. Olefinic Compounds .................................................................................................... 1465 1. Hydrogenation to Alkanes .................................................................................... 1465 2. Hydrodimerization ................................................................................................ 1465 B. Acetylenic Compounds ................................................................................................ 1468 C. Aromatic Compounds .................................................................................................. 1469 D. Carbonyl Compounds .................................................................................................. 1469 1. Hydrogenation to Alcohols ................................................................................... 1469 2. Hydrodimerization to Pinacols ..............................................................................1471 3. Cross-Dimerization with Unsaturated Systems .................................................... 1472 E. vic-Oxygen Compounds ...............................................................................................1474 F. Nitrogen Compounds ....................................................................................................1474 G. Sulfur Compounds ........................................................................................................1475 H. Halogen Compounds ....................................................................................................1475 1. Cyclic Halides ........................................................................................................1475 2. Acyclic Polyhaloalkanes ....................................................................................... 1477 3. Acyclic Halo-Olefins............................................................................................. 1477 4. Intermolecular Dimerization ................................................................................ 1478 5. Intramolecular Cyclization ................................................................................... 1478 I. Other Heteroatom Compounds .................................................................................... 1479 J. Organometallic Compounds ........................................................................................ 1479 IV. Stereoselective and Stereospecific Electrooxidation ........................................................... 1479 A. Carboxylic Acids ......................................................................................................... 1479 1. Kolbe-Type Reaction ............................................................................................. 1479 2. Non-Kolbe-Type Reaction..................................................................................... 1480 B. Acetoxylation ............................................................................................................... 1480 1. Olefins ................................................................................................................... 1480 2. Benzyl and Allyl Compounds ............................................................................... 1480 3. Cyclopolyenes ....................................................................................................... 1481

1461

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C. Methoxylation and Hydroxylation ............................................................................... 1481 1. Olefins ............................................................................................................................ 1481 2. Aromatics .............................................................................................................. 1481 3. Others .................................................................................................................... 1481 D. Fluorination ................................................................................................................. 1482 E. Acetamidation.............................................................................................................. 1483 1. Haloacetamidation ................................................................................................ 1483 2. Thioacetamidation ................................................................................................ 1483 F. Other Addition Reactions to Olefins............................................................................ 1483 1. Halogenation ......................................................................................................... 1483 2. Selenoalkoxylation, Thioalkoxylation, Thiohydroxylation, and Disulfenylation.... 1484 G. Intermolecular and Intramolecular Reactions of Olefins and/or Aromatics ............... 1484 1. Dimerization ......................................................................................................... 1484 2. Cross-Coupling ..................................................................................................... 1485 H. Active Hydrogen Compounds (Carbon Acids) ............................................................ 1485 I. Heteroatom Compounds .............................................................................................. 1486 1. Nitrogen Compounds ............................................................................................ 1486 2. Sulfur Compounds ................................................................................................ 1486 3. Silicon Compounds ............................................................................................... 1487 J. Organometallic Compounds ........................................................................................ 1487 V. Asymmetric Reduction and Oxidation................................................................................ 1487 A. Reaction of Chiral Compounds ................................................................................... 1487 1. Reduction of Halides............................................................................................. 1487 2. Reduction of Alcohols and Amines ...................................................................... 1487 3. Reduction of Onium Compounds ......................................................................... 1488 4. Oxidation of Carboxylic Acids and Alcohols ....................................................... 1488 5. Oxidation of Heteroatom Compounds .................................................................. 1488 B. Asymmetric Synthesis ................................................................................................. 1489 1. Use of Chiral Electrode Adsorbants ..................................................................... 1489 2. Use of Chiral Supporting Electrolytes .................................................................. 1489 3. Use of Chiral Media.............................................................................................. 1490 4. Use of Chiral Oxidant ........................................................................................... 1490 5. Use of Chiral Modified Electrodes ....................................................................... 1490 6. Intramolecular Asymmetric Induction ................................................................. 1491 7. New Asymmetry Induction Methods.................................................................... 1497 C. Electrochemical Kinetic Resolution ............................................................................ 1498 VI. Electrochemical Stereoisomerization.................................................................................. 1499 A. Cis-to-Trans Isomerization .......................................................................................... 1499 B. Tub-to-Chair Isomerization ......................................................................................... 1499 C. A-to-B Form Isomerization in Bianthrones................................................................. 1500 VII. Stereoisomeric Effects on Thermodynamics and Kinetics of Electrode Processes ........... 1500 A. Cathodic Processes ...................................................................................................... 1500 1. Hydrocarbons ........................................................................................................ 1500 2. Halogen Compounds ............................................................................................. 1500 3. Nitrogen Compounds ............................................................................................ 1500 4. Organometallic Compounds ................................................................................. 1501 B. Anodic Processes ......................................................................................................... 1501 1. Hydrocarbons ........................................................................................................ 1501 2. Alcohols ................................................................................................................ 1501 3. Other Heteroatom Compounds ............................................................................. 1501 References .................................................................................................................................... 1501

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I. INTRODUCTION An electrochemical reaction is a unique oxidation–reduction that takes place in a specific heterogeneous interface between an electrode and an electrolytic solution, as is described in detail in the following text. Therefore, the stereochemical features of electrochemical reactions must be also unique. However, generally, it is not easy to rationalize simply and clearly a stereochemical course involved in a given electrochemical reaction on the basis of experimental results and theoretical considerations, since the manner of stereocontrol is very complicated in general. Thus, despite a number of rules, it is also not easy to predict the stereochemical results in advance. However, nowadays, it is not so difficult to optimize stereoselective synthesis empirically. In 1928, Clemo and Smith investigated the cathodic reduction of p-dimethylaminobenzaldehyde in aqueous sulfuric acid and separated one of the two diastereomeric glycols as a product [1]. This is the pioneering work of the stereochemical study on organic electrode processes. Later, along with developments in organic synthesis, a number of original papers dealing with stereochemical aspects of electrode reactions have been published, and some general [2–7] and topical [8–15] reviews should be very helpful for this type of investigation. Although papers published after 1980 are mainly cited as references in this chapter, readers can access those before 1980 by referring to the previous third and fourth editions of this book [2,3].

II.

STEREOCONTROL IN ELECTROCHEMICAL REACTIONS

A. STERIC AND POLAR FACTORS Substrate and/or intermediate species adsorb on an electrode surface and orient themselves so that their least hindered sides face the electrode unless there is another effect, such as a polar one. This may be the simplest steric factor governing the stereochemistry of reactions. There may also be more complicated steric effects that result in conformational or configurational change of the species, since the electrode on which the species adsorbed strongly may behave as if it were a very bulky substituent. For example, visualize a substituted cyclohexane molecule (Figure 38.1). The substituent must be preferentially at the equatorial position in a free state (1a), but it may be at the axial position in an adsorbing state (1b), if the electrode effectively gives more steric hindrance than the substituent. In addition to the steric interaction of adsorbed species with the electrode, the species can sterically interact with the other species adsorbed on the electrode surface. An electrode interface has a layered structure in which a nonuniform electric field (potential slope) is generated by polarization of the electrode. An extremely strong electric field of around 108 V cm−1 in the most inner layer, the so-called electron-transfer layer, which is very thin, 10 Å or less, might cause a variety of polar effects. Since not only the electron-transfer step but also adsorption and some of the chemical steps involved in an electrolytic reaction take place in the electron-transfer layer, the electrochemical reaction should be strongly influenced by polar factors. The orientation of polar adsorbed species, such as ions and dipoles, is electrostatically influenced, and consequently, the stereochemistry of their reactions is also controlled by this kind of electrostatic factor. Electrode R R

1a

FIgURE 38.1

1b

Steric interaction of adsorbed substituted cyclohexane with an electrode.

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Organic Electrochemistry

A conformational or configurational change may also occur at the electrode interface. For example, when a planar radical (2a) receives one electron from a cathode to give the corresponding anion, the original configuration of the radical may be inverted by such a strong electric field to give the inverted anion (2b), as shown in Equation 38.1. Cathode – e–

C

2a Planar radical

B.

(38.1)

C

2b Inverted anion

CONFORMATIONAL AND CONFIGURATIONAL STABILITIES

The stereochemistry of reactions is controlled by conformational and configurational stabilities of intermediate species in the transition state. Therefore, the stereochemical course of the reactions must be discussed on the basis of the intermediates and reaction mechanism that have been presumably clarified by electrochemically and organic chemically reasonable ways. However, on the contrary, it is also likely that a stereochemical result makes the mechanism more clear. In this sense, a stereochemical study can be useful as one of methods for analyzing the mechanism. Anyway, it should be noted that the conformational and configurational stabilities of the species in an electrode interface may be quite different from those in free states, owing to the peculiar steric and polar factors. In a given situation, the conformation or configuration at the electrode interface may be the reverse of that in the free state. The relation between the rate of each reaction step and the steric stability of the intermediate species must also be considered to rationalize the stereochemical results. Consider the electrochemical reduction (Equation 38.2) of an alkyl halide (R–X) to the corresponding alkane (R–H) via an anion radical [(R–X)• −], radical (R•), and anion (R−) mechanism (Electrochemical-ChemicalElectrochemical-Chemical reactions; ECEC). The order of configurational stability may be R–H > (R–X)• − > R− >R• in general, but the stability of R• is not always lower than that of R− when the adsorption of R− is weakened by electrostatic repulsion by the cathode. The situation is more complicated when the stereoisomerization of the intermediates must be considered. For instance, if the reaction rate of an unstable intermediate is high enough, its isomerization can actually be neglected. In contrast, the isomerization of other intermediates with very small rates cannot be neglected even though they may have considerable stability.

R–X

e–

[R–X]

–X–

–

R

e–

R–

H+

R–H

(38.2)

Stereoisomerization [R´–X]

C.

–

–X–

R´

e–

R´–

H+

R´–H

THERMODYNAMIC AND KINETIC CONTROLS

It is important in a mechanistic discussion of stereochemical results to know whether the stereochemistry of products is kinetically or thermodynamically controlled. The thermodynamic equilibrium ratio of stereoisomeric products should be considered. The computational molecular calculation is commonly used to obtain the theoretical thermodynamic equilibrium ratio.

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D.

STEREOSELECTIVE SYNTHESIS

It has been well known that stereochemical results are affected by a variety of conditions, such as electrode material, electrode potential, current density, supporting electrolyte, solvent, additive, and temperature. Modification of electrodes is also useful for changing stereochemical results. Therefore, in this chapter, the theoretical and mechanistic problems of the stereochemistry of electrochemical reactions are discussed, as well as stereoselective synthesis.

III.

STEREOSELECTIVE AND STEREOSPECIFIC ELECTROREDUCTION

A. OLEFINIC COMPOUNDS 1. Hydrogenation to Alkanes Since olefin molecules can have stereoisomers (cis and trans; Z and E), the interesting stereochemistry, which is not only stereoselectivity but also stereospecificity, of their hydrogenation should be observed when the products also have stereoisomers (d,l and meso; threo and erythro). However, only in a few works could stereospecificity be observed. Dimethylmaleic and dimethylfumaric acids are reduced at a mercury cathode to form d,l and meso α,α′-dimethylsuccinic acids, respectively. This is direct evidence that the hydrogenation proceeds through trans addition of hydrogen. A remarkable stereoselective hydrogenation of a C=C double bond was reported as follows: 2,3-Diphenylindenone undergoes a selective cis addition in 0.25 M sulfuric acid, while a trans addition was found at higher pH values (Equation 38.3) [16]. This suggests that the hydrogenation proceeds through the cis addition in a strongly acidic medium but with trans addition in approximately neutral media. Ph

Ph

Ph 2e–, +2H+

2e–, +2H+ Ph O

Ph

pH = 4.7–9.3 O

Ph

0.25 M H2SO4

(38.3)

O

Utley and coworkers [17,18] also reported that the stereoisomeric ratio (cis/trans) of 1,4-disubstituted cyclohexanes formed by the hydrogenation of the corresponding activated cyclohexenes at a mercury electrode depended on solvents and proton sources, and discussed the stereochemical mechanism in detail. On the other hand, according to Lessard and coworkers [19], the hydrogenation at a Raney nickel electrode always provides the trans-isomers in excess (see Chapter 44). The hydrogenation of conjugated dienes can also have stereochemical features; however, only few mechanistically valuable studies have been reported [20–22]. 2. Hydrodimerization a. Intermolecular Dimerization The intermolecular hydrodimerization of olefins gives stereoisomeric (d,l and meso) products. Moreover, the stereospecificity can be discussed by using stereoisomeric (cis and trans; Z and E) olefins. Baizer and coworkers established the most brilliant industrial electroorganic synthesis of the hydrodimerization of acrylonitrile to adiponitrile. They extended this hydrodimerization to a variety of activated olefins [23–25] and paid only little attention to the stereochemistry of products. However, their stereochemical data were not enough to discuss the stereochemical course of the reaction. Utley and coworkers [26–28] have reported stereochemical data of hydrodimers derived from a variety of cinnamic acid esters with chiral alcohol components.

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Simonet and coworkers [29,30] also reported the stereochemistry of hydrodimers derived from activated olefins in detail. The d,l-dimers were predominant in most cases; however, the remarkable substituent effect was not enough explained. Some fundamental studies were made earlier. It was found that the hydrodimerization of an α,βunsaturated ketone (enone) was at least partially stereospecific, with the d,l dimer formed from the trans-enone and mainly the meso dimer formed from the cis-enone. Kanetsuna and Nonaka [31] also examined both the stereoselectivity and stereospecificity of the hydrodimerization in detail. It was indicated that either the reaction mechanism of the stereochemical course is changed by the nature (protic or aprotic) of the reaction media. Particularly, the latter is greatly influenced by the manner of adsorption of an anion radical intermediate on the cathode. It was also postulated to explain the experimental results that the anion radical adsorbs at the β-position, where the radical π-electron is delocalized and the adsorption is affected by the nature of the media. Lund, Utley, Simonet, and coworkers reported that the reductive cyclodimerization of activated olefins in aprotic media affords the trans-1, 2-disubstituted cyclobutanes as shown in Equation 38.4 [32–34]. EWG 2

(38.4)

EWG EWG

The reductive cross-coupling between acrylates with butadiene provided the corresponding transunsaturated carboxylic acid esters almost predominantly (Equation 38.5) [35].

H2C — CH CH

H3C CH2 +

C

CHCOOR

H3C

2e–

C— —C

DMF

H3C

CH3 CH2 C CH2COOR CH3

(38.5)

The hydrodimerization of nonactivated olefins at a platinum electrode affords the corresponding d,l-isomeric products in excess (>80%) [36]. b. Intramolecular Cyclization Since the development of the hydrodimerization of activated olefins in aqueous solutions containing quaternary ammonium salts, some workers have paid attention to the stereochemistry of reductive intramolecular cyclization of diolefins as well as the intermolecular dimerization of monoolefins. The reaction is presented in Equation 38.6, where the trans-cyclized products are always predominant [37,38]. These results seem to be interesting from the aspect of the stereoselective synthesis of cyclic compounds, although the stereochemical course of each reaction has not been made clear in most cases because of uncertainty about the stereochemistry of the starting diolefins. CHCOOR n

CHCOOR n = 1–4

2e–, +2H+

CH2COOR n

(38.6)

CH2COOR

37–87% (cis/trans = 0–0.93)

It is noticeable that the cis–trans ratio of the reaction products is influenced strongly by proton donor and cathode material [38]. The ratio and the yield increase when a stronger proton donor is used. This may suggest that the protonation of anionic intermediate species is an important key step in either cyclization or stereocontrol. The fast protonation of the configurationally unstable anionic species results in the predominant formation of the less stable cis products.

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c. Cross-Coupling and Cyclization Reductive cross-coupling of olefins with anionic and radical species formed by electroreduction of other molecules is also important from both the synthetic and mechanistic aspects (see also Chapter 17). d. Acylation with Acetic Anhydride Shono and coworkers [39] reported that the reduction of ethyl 1-cyclohexenecarboxylate in an anhydrous acetonitrile solution in the presence of acetic anhydride afforded the trans product with moderate stereoselectivity, as in Equation 38.7. COOEt

+

Ac2O

COOEt

2e–

(38.7) Ac 74% (cis/trans = 0.32)

e. Carboxylation with Carbon Dioxide It is well known that olefins, when reduced in the presence of carbon dioxide, are dicarboxylated in high yields, but poor stereoselectivity was found. f. Cross-Coupling with Carbonyl Compounds The electroreductive cross-coupling of olefins with carbonyl compounds, which are initially reduced to radical or anion intermediates to attack the olefins, is synthetically useful. The stereochemistry of this type of intermolecular reaction was first examined [40–42], and later that of intramolecular cyclization was more intensively investigated [43–45]. Kanetsuna and Nonaka [40,41] examined in detail the stereochemistry of the cross-coupling products of crotonic acid derivatives with dissymmetrical carbonyl compounds (Equation 38.8). The cis–trans ratio of the cyclized products varied in a wide range of 2.2–0.22 depending on the molecular structure of the starting compounds and electrolytic conditions.

– –

X Me CH — —C

Me +

Y

2e–, +2H+ O

Me CH

(38.8)

Me C OH

Z

Z threo/erythro

X = H, COOH, CN Y = H, COOH, COOMe, CN

Hydrolysislactonization

Z = H, Et Me Me

O O

Z cis/trans = 2.2–0.22

g. Cyclization of Olefinic Carbonyl Compounds The stereochemistry of intramolecular cyclization of olefinic carbonyl compounds has been extensively studied [43,44,46]. Particularly in the case of unactivated olefins with carbonyl functions, which are reduced more easily than the C=C double bonds, the stereoselectivity is so high that only a single isomer is formed, as in Equation 38.9 [43,47].

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Organic Electrochemistry

CH2

Me

2e–, +2H+

CH(CH2)n CR

HO R

O

(38.9)

n-2

35–98% (single isomer)

An interesting extension is the twofold cyclization of nonconjugated dienones to form bicyclic alcohols in high yield and with high stereoselectivity (Equation 38.10) [48]. (+)-N,N′-dimethylquininium (DMQ) tetrafluoroborate causes an anodically shifted catalytic current and is essential for the double cyclization. It is proposed that it is adsorbed at the cathode and facilitates reduction and exo-trigcyclization by hydrogen bonding. O–

O

OH H+, +e–

e–

(38.10)

H+

DMF/i-PrOH DMQBF4

H

H

65% H HO H DMQ =

MeO

N+ +

Me

N Me

Compared with unactivated olefins, activated olefins with carbonyl groups are reductively cyclized with somewhat little lower stereoselectivity [44,49–51]. The reductive cyclization of unsaturated carbonyl compounds can also be achieved by using chemical reducing reagents, such as Al-amalgam, MgTiCl4, and Na–NH3, but in general, not only chemical yields of cyclized products but also stereoselectivity is lower than those of the electroreductive cyclization. Pattenden and Robertson [52] reported a very interesting cathodic reduction of a variety of terminal allenic ketones that were cyclized highly stereoselectively through the exo mode, producing five-membered rings. It is particularly notable that the products formed in the electrochemical and chemical (sodium naphthalenide) reductions are two structural isomers with the same configuration (Equation 38.11). OH 2e–, +2H+ H

O

OH C10H8–Na+

(38.11) H

Electroreduction of nonconjugated enones and ynones containing a sulfur or a nitrogen atom in the connecting chain resulted in regio- and stereoselective cyclization to afford heterocyclic alcohols in moderate to good yields [53].

B. ACETYLENIC COMPOUNDS The stereochemical feature of electrolytic hydrogenation of acetylenes to the corresponding olefins changes drastically with electrolytic conditions, such as supporting electrolyte, solvent, and cathode material.

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Kita and Nakajima [54] found that the cis–trans ratio of 6.3 of 2-butene formed from 2-butyne at a palladium black cathode was not influenced by neither cathode potential nor pressure. On the other hand, Nonaka and coworkers [55] examined the cis–trans ratios of dimethyl-2-butenedioate formed from dimethyl acetylene dicarboxylate on a palladium black cathode and catalytic hydrogenation at various pH levels and found that an increase in pH resulted in a decrease in the ratio in both the electrolytic and catalytic hydrogenations, but the former was less strongly influenced by pH. They also performed similar comparative studies with olefins and ketones from a stereochemical aspect. An interesting electroreductive cyclization of acetylenic ketones was reported, as shown in Equation 38.12 [56]. The cyclized products were formed in good yields with moderate stereoselectivity. Perichon and coworkers [57] also reported that the reductive cross-coupling of 1,3-dynes with carbon dioxide gave exclusively the corresponding cis-addition products in the presence of a Ni(II) complex, as shown in Equation 38.13. R O

R

2e–, +2H+

(38.12)

Et4NOTs/DMF

OH 53–60% (E/Z = 0.5) R1

R1

2e–, +2H+

R2 + CO2

H

(38.13)

R2

Ni(II)

COOH

C. AROMATIC COMPOUNDS Aromatic hydrocarbon nuclei can be hydrogenated partially or completely under various electrolytic conditions. Hydrogenation of phthalic acid to dihydrophthalic acid is industrially important. Phthalic acid is efficiently hydrogenated at high hydrogen overpotential cathodes, such as mercury and lead in acidic solutions, to give almost exclusively trans-dihydrophthalic acids [58]. In contrast, o-methylphenol is stereoselectively hydrogenated on an Rh/C cathode to give cis-2-methylcyclohexanol (42–83% yields [cis/trans ratio=2.3–6.7]) as the major isomer [59].

D.

CARBONYL COMPOUNDS

The stereochemistry of the cathodic reaction of carbonyl compounds has been intensively studied from both synthetic and mechanistic aspects. In this section, the reaction is discussed in three parts: hydrogenation to alcohols, hydrodimerization to pinacols, and cross-dimerization with unsaturated compounds. 1. Hydrogenation to Alcohols a. Acyclic Carbonyl Compounds As shown in Equation 38.14, the reduction of 1,2-diphenyl-1-propanone at an Hg cathode in aq. EtOH and pH 8 provides the erythro alcohol as the major diastereomer (erythro:threo=5–1.4:1) [60]. This selectivity is in accordance with a protonation of the intermediate anion, formed in an Electrochemical-Chemical-Electrochemical (ECE) sequence, from the least hindered side. OH

O Ph

Ph

e–, +H+ e–

CH3

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Ph

C– Ph

HO CH3

CH3

Ph H H+

H Ph

H

Ph H H

OH CH3 Ph

(38.14)

1470

Organic Electrochemistry

Later, Nonaka and coworkers [61] also studied the cathodic reduction of racemic ketones and found that the stereochemistry of the products was affected greatly by R1 and R2 substituents of the starting ketones and the pH of the electrolytic solutions used as shown in Equation 38.15. It is noticeable that when R1=cyclo-Hex and R2=Me, the anti-Cram isomer was formed as a major product. This suggests that stereochemical interaction between the cathode and intermediates contributes to stereocontrol in the reaction. 2e–, +2H+

Me O

H H R1–C –C –R2 – – – –

– – – –

H R1– C – C – R2

(38.15)

Me OH R1 = Ph, R2 = Ph: 8–91% (threo/erythro = 0.2–0.7) R1 = Ph, R2 = Me: 3–82% (threo/erythro = 1.3–1.4) R1 = c-Hex, R2 = Me: 3–7% (threo/erythro = 0.5–0.9)

A similar study using another type of starting ketones with a chiral center at the β-position affords diastereomers in a ratio different from that obtained by LiAlH4 reduction [62]. b. Cyclic Carbonyl Compounds Most of the stereochemical studies on the hydrogenation of cyclic carbonyl compounds to the cyclic alcohols have been performed using substituted cyclic cyclohexanones and cyclopentanones [55,63–67]; the stereochemistry of the cyclic alcohols formed is discussed by examining the cis– trans geometry and/or ax-eq conformation of the hydroxy groups. A t-butyl group substituted on a cyclohexane ring is always at the equatorial position. Therefore, for instance, it can be reasonably explained that the less stable cis-4-t-butylcyclohexanol with an ax-OH is formed predominantly (98% cis form) in electrolytic solutions containing bulky cations, such as Mg2+ and Zn2+, which make tight ion pairs with carbanions as shown in Equation 38.16 [64,65]. Cathode O

OH e–, +H+

OH e–

–

Diffusion OH H cis (ax-OH) form

OH – + MgCl

H+

+ MgCl

(38.16)

– OH

Ion pair

4-t-Butylcyclohexanone is reduced at a Pt-black cathode to 4-t-butylcyclohexanol with a cis/trans ratio of 0.9 in acidic and 1.0 in basic media. At an Hg cathode, the corresponding ratios are 0.2 and 0.4 [55]. For 2-methylcyclohexanone, the cis/trans ratio of the formed 2-methylcyclohexanol depends on the electrode material, current density, cosolvents, and pH of the electrolyte. In a basic medium, the best yields were obtained with a cis/trans ratio of 35/65 [66]. For 4-t-butylcyclohexanone and 3,3,5-trimethylcyclohexanone at a Pt-cathode in LiCl, EtOH, and HMPA, a good yield of the alcohol with a eq/ax ratio of 95/5 was obtained, while at a Pb-cathode in NaOAc, HOAc, and MeOH, a lower yield and a eq/ax ratio of 44/56 was observed [68]. A series of cyclic ketones was reduced in iso-propanol and in H2SO 4, H 2O, and MeOH. It was found that in iso-propanol, the ratio of diastereomers was approximately equal to their relative thermodynamic stabilities, while in an acidic medium, the less stable epimeric alcohols were formed to a higher extent. This result was taken as an indication that in acidic medium, an adsorbed

© 2016 by Taylor & Francis Group, LLC

1471

Stereochemistry of Organic Electrode Processes

intermediate is rapidly protonated, while in the poor proton-donating iso-propanol, this intermediate can diffuse into solution and equilibrate this way [63]. 2. Hydrodimerization to Pinacols a. Intermolecular Dimerization Many papers on hydrodimerization of aromatic carbonyl compounds have appeared indicating the importance of this reaction. The rac/meso ratio for the pinacolization of acetophenone in aqueous ethanol ranges between 0.9 and 1.4 in acidic medium and between 2.5 and 3.2 in basic medium. The diastereoselectivity is independent of the cathode material, Hg, Sn, or Cu. Electrolysis conditions such as current density, potential, or current controlled electrolysis also do not influence the diastereoselectivity. The same holds for propiophenone. For benzaldehyde the rac/meso ratio is 1.1–1.2 in acidic as well as in basic media [69,70]. In the presence of adsorbable ions such as I−, Et4N+, the rac/meso ratio for benzaldehyde decreases to 0.5 [71]. In dry DMF/LiClO 4, the rac/meso ratio for acetophenone increases up to 12.5 and 19 [72,73]. It was proposed that in acidic solution, two ketyl radicals couple and in alkaline solution, a ketyl radical and a radical anion couple. There is one transition state 3a leading to meso product and two 3b, 3c forming rac product as shown in Figure 38.2. 3a minimizes steric interactions, while in 3b, 3c, hydrogen bonding is involved. The latter is more or less pronounced in acidic and basic protic media. In aprotic media with complexing Li+, ion bridging is maximized, leading to high rac selectivity. Some α,β-unsaturated carbonyl compounds are stereoselectively hydrodimerized. For example, cathodic reduction of optically active compound 4 provided three isomers, of which 5 was major over a wide range of pH as shown in Equation 38.17 [74,75].

e–, +H+

(38.17)

1/2

O

HO OH 5 cis-threo-cis (major)

4

Although the stereochemistry of hydrodimers is greatly affected by the structure of the starting carbonyl compounds and electrolytic conditions, the less stable d,l-isomers are very rarely formed in d,l-meso ratios higher than 1.5 in the hydrodimerization of simple aromatic carbonyl compounds. Rusling and Zuman [76] reported that 4-pyridylcarboxaldehyde gave the hydrodimer in the d,l-meso ratio of 1.92 in a strongly alkaline solution. Furthermore, Smith and Utley [77] found benzaldehyde was hydrodimerized at a mercury cathode in dry DMF solution of Bu4NBF4 to give the d,l-meso ratio of 14; the addition of β-cyclodextrin drastically decreased the ratio to 0.33. Notably, the meso/rac ratio in the electro-hydrodimerization of acetophenone in weakly acidic solution could be influenced by ultrasonication [78]. H

H

O

OH H3C

Ph

H3C

Ph

CH3

Ph

O O H CH3

H

O

Ph CH3

Ph

OH

Ph

CH3

3a

3b

3c

 

FIgURE 38.2 Transition state of hydrodimerization of phenyl methyl ketone in acidic protic media.

© 2016 by Taylor & Francis Group, LLC

1472

Organic Electrochemistry

b. Intramolecular Cyclization Some aromatic diketones have been stereoselectively cyclized under various electrolysis conditions, which together with the substrate structure strongly influence the stereochemistry of the formed cyclic diol. Reductive cyclization of 1,8-diaroylnaphthalenes led to trans-diols; 2,2′-diaroylbiphenyls and α,ω-diaroylalkanes yielded cis-diols with different stereoselectivities, depending on substrate structure and electrolysis conditions (pH and cosolvent), as in Equation 38.18 [79,80]. 2-Substituted 1,3-diphenyl-1,3-propanediones provided in high yield and cis-selectivity of 3-methyl substituted 1,2-diphenyl-1,2-cyclopropanediols [81]. 1-Acyl-9-benzoylnaphthalenes were cyclized to 1,2-acenaphthenediols in 50–100% yield and a cis/trans ratio ranging from 100:0 (preferably in an acidic medium) to 0:100 (preferably in a basic medium), which is also dependent on the kind of acyl group, supporting electrolyte, cosolvent, and cathode potential. The preferable formation of the cis-diol in acidic medium is rationalized as due to cyclization in a radical–radical coupling, while in alkaline medium, the trans-diol is formed by intramolecular nucleophilic addition of the radical anion to the carbonyl group [82].

– –

– –

R1 – C – R3 – C – R2 O

O

2e–, +2H+

R1 HO

R1

R2

R3

Yield (%)

C6H5 C6H5 C6H5 4-CH3OC6H4 4-HOC6H4 4-CH3OC6H4

C6H5 C6H5 Me 4-CH3OC6H4 4-HOC6H4 4-CH3OC6H4

CHMe CMe2 CHMe (CH2)3 (CH2)3 (CH2)4

80 80 80 82 89 64

R3

R2

(38.18)

OH cis/trans-OH cis only cis only 1 26 0.9 cis only

Stoichiometric amounts of ytterbium(II), generated by reduction of Yb(III), supports the stereospecific coupling of 1,3-dibenzoylpropane to cis-cyclopentane-1,2-diol. However, Yb(III) remains bounded to the pinacol and cannot be released to act as a catalyst. This leads to a loss of stereoselectivity in the course of the reaction [83]. 3. Cross-Dimerization with Unsaturated Systems In most cases of the cross-dimerization between carbonyl and unsaturated groups in Section III.A.2, there has been an ambiguous point in which the carbonyl or unsaturated group is first reduced to give a reactive intermediate species. Hence, the stereochemical course of reaction has not been sufficiently discussed. In this section, the cross-dimerization, in which it has been mechanistically clarified that the carbonyl groups are reduced to give the corresponding radical or anion intermediate species that react with the unsaturated bonds such as C=C, C=N, and C≡N, and the stereochemical course are discussed in detail. a. Cross-Dimerization with C=C Bonds Cathodic coupling of ketones with allylic alcohols was reported by Shono and coworkers [84] to take place at a carbon fiber cathode with high diastereoselectivities due to interaction (hydrogen bonding) between two hydroxyl groups of the starting alcohols and anion radical and radical species derived from the ketone. It was also found that the reduction of nonconjugated aromatic ketones at metal cathodes gave the corresponding cis-isomers of intramolecularly cyclized products with high stereoselectivity [84], similar to the reduction shown in Equation 38.9. Reduction of nonconjugated aromatic ketones at metal cathodes (e.g., Sn, Cu, Ag, Pd, and Zn) gave the cis-isomers (cis-H/OH) of cyclized products in high diastereoselectivity. The electroreduction of

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1473

Stereochemistry of Organic Electrode Processes

5-phenylpentan-2-one led to 70% of an exclusively cis-hexahydronaphthalene using an Sn cathode and tetraalkylammonium salts in i-PrOH (Equation 38.19). A variety of new bi- and polycyclic tertiary alcohols were prepared in this way. Mechanistic studies point to a reduction of the carbonyl group to a radical anion that attacks the aromatic ring in such a way that the negatively charged oxygen atom avoids electronic repulsion with the π-electrons of the aromatic ring, leading to the product with H and OH being cis to each other. Reduction with Na in HMPA-THF gave the same cyclized product, however, in lower yield than in the electroreduction [45,85]. O

H

Me OH

Sn cathode

(38.19)

Et4NOTs/i-PrOH 70%

b. Cross-Dimerization with C=N Bonds Reductive cross-dimerization has been established with ketones and O-methoximes upon reduction in iso-propanol with an Sn cathode as a convenient route to β-amino alcohols; diastereoselectivities of trans-cis ratio up to 95:5 were obtained. A chiral ligand was in this way obtained from the coupling of (−)-menthone with O-methyl acetaldoxime. Similarly ketones could be coupled to hydrazones and nitrones. Also intramolecular couplings were achieved with good yields and diastereoselectivity (Equation 38.20) [86]. OMe

OMe N

O

HO HN

Sn cathode

H

H

(38.20)

Et4NOTs/i-PrOH n

n

n = 1: 65%, trans/cis = 95/5 n = 2: 75%, trans/cis = 85/15

Regioselective intramolecular coupling is also found in the cathodic reduction of oxoalkyl pyridinium salts to yield polycyclic pyrrolidines and piperidines (Equation 38.21) [87,88]. The reaction proceeds presumably via reduction of the hydrogen-bonded carbonyl group to a nucleophilic hydroxyalkyl radical, that undergoes addition to the electrophilic C=N double bond of the pyridine. Br– +

O

4e–, +3H+

H

HO

H +

H N

1 M H2SO4

HO H

(38.21)

N

N 73% 13:1

c. Cross-Dimerization with C≡N Bonds Electroreduction of γ- and δ-cyano ketones in iso-PrOH with an Sn cathode gave α-hydroxyketones as cyclization products with good diastereoselectivity as shown in Equation 38.22 [89]. The reaction has been used as key step for the synthesis of, for example, guaiazulene, triquinanes, and dihydrojasmone. Similarly, the corresponding intermolecular couplings were realized. CN n m

© 2016 by Taylor & Francis Group, LLC

O

H 2e–, +2H+ –NH3

n m

OH O

(38.22)

1474

E.

Organic Electrochemistry VIC-OXYGEN

COMPOUNDS

Cathodic reduction of benzil in acidic and neutral solutions, followed by dimethylation, provides cis-stilbenediol dimethyl ether predominantly [90]. 1,3-Dimethanesulfonates undergo electroreductive cyclization in DMF containing ammonium salts to give stereoselectively the corresponding trans-cyclopropanes in a good yield, as in Equation 38.23 [91]. OMs n-C12H25

2e– n-C12H25

(38.23)

2 MsO–

+

OMs 76% (trans/cis = 3/1)

A stereoselective one-electron reductive elimination reaction of cyclic sulfates of 1,2-diols proceeds to provide the trans-olefins selectively [92]. For instance, the reductive elimination of the cyclic sulfate of 1,2-butanediol led to trans-2-butene selectively. Aromatic vicinal dioxalates underwent fragmentation and elimination on cathodic reduction to give alkenes. meso-(EtO2COCH(C6H5))2 gave 80% trans-stilbene [93].

F. NITROGEN COMPOUNDS Electrochemical reduction of camphoroxime and norcamphoroxime at an Hg cathode proceeds with a high degree of stereoselectivity to give products of opposite stereochemistry to those formed in dissolving metal (Na-alcohol) reduction of the oximes. The electrolyses are proposed to proceed by a kinetically controlled attack by the electrode on each oxime from the less hindered side as shown in Equation 38.24 [94]. In contrast, the corresponding N-phenyl imines provide products of the same stereochemistry as those isolated from a dissolving metal reduction. Cyclic voltammetry and polarographic data point to RH − and R2− intermediates in this case that are protonated from the least hindered side [95]. R

R

R

R

R

R

Hg cathode LiCl/CH3OH, H2O

R

NOH

(38.24)

+

NH2 R

R

NH2

50–70% R = H, 0:100 R = CH3, 99:1

An acyclic vic-diimine was hydrogenated at a mercury cathode to give a mixture of the d,l-and meso-diamines in 70% yield [96]. Kise and coworkers found efficient intramolecular hydrodimerization of aromatic diimines using a lead cathode providing 1,4-diazacrown ethers, as in Equation 38.25 [97]. The stereoselectivity greatly depended on the ring size of the products and solvents used. Ar

N

O

Ar

N

O

n

2e–, +2H+

Ar

NH

O

DMF or DMF/THF

Ar

NH

O

n

n = 0, Ar = Ph: cis/trans = 0 n = 1, Ar = Ph, 1-furyl: cis/trans = 0.053 n = 2, Ar = Ph: cis/trans = 0.33 n = 5, Ar = Ph: cis/trans = 0.67–1.5

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(38.25)

1475

Stereochemistry of Organic Electrode Processes

Cathodic reduction of an imidazolinone derivative at Hg cathode provided hydrogenated products with moderate diastereoselectivity. The diastereomeric ratio depended on pH of the electrolytic solution (Equation 38.26) [98]. COOH

i

N

Me

N H

N

Pr –

+

2e , +2H

COOH H N

iPr

N H

O

N

O

Me

(38.26)

dr = 0.77 (pH = 1.2) = 1.86 (pH = 10.9)

G. SULFUR COMPOUNDS Sulfur compounds with oxygen functions at the β-positions are reductively desulfurized to give the corresponding trans-olefins, regardless of the original stereochemistry of threo/erythro, as in Equation 38.27 [99,100]. The reaction is highly stereoselective but not stereospecific. R1

R2

2e–

H

H

R2

(38.27)

OR

PhS

R1

(PhSO2)

Some ring-opening reactions of cyclic sulfur compounds to olefins proceed stereoselectively via desulfurization. 1,2-Diphenylthiirene dioxide is reduced to eliminate sulfur dioxide, giving transstilbene in 30% yield [101]. The reductive coupling of dithiobenzoate esters results in the predominant formation of Z-1,2dialkylthio-1,2-diphenylethylenes in good yields when alkyl halides are added after completion of the electrolysis, as shown in Equation 38.28 [102]. Ph

C

SR1

2e–

Ph

R2I

R1S

Ph

(38.28)

S

SR2

Cathodically initiated Michael addition of thiols to levoglucosenone provides threo or erythro addition products selectively depending on the amount of charge passed, as in Equation 38.29 [103]. O

O

0.05 F

O

RS

(38.29)

O O

O threo

H. HALOGEN COMPOUNDS A variety of stereoselective and stereospecific reductions of halogen compounds have been reported. 1. Cyclic Halides a. Cyclic Monohalides Only a few reports on the stereochemistry of reduction of monohalides, except chiral halides (see Section V), have appeared so far. Tallec and coworkers [104] found that the reduction of cis-and

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1476

Organic Electrochemistry

trans-cyclopropyl bromides proceeds stereospecifically in a tetrabutylammonium salt solution. On the other hand, in an ammonium (NH4+) salt solution, the opposite stereochemistry of the products is observed. b. Monocyclic gem-Dihalides 1-Substituted 2,2-dibromocyclopropanes are stereoselectively reduced to the corresponding diastereomeric monobromides (a and b) in good yields, as shown in Equation 38.30 [105]. Although the stereochemistry of the monobromides formed is strongly influenced by the substituents (R1 and R2) of the starting substrate and electrolytic conditions, such as pH, supporting electrolyte in aqueous ethanol, and cathode potential, erythro-type isomer (a) is always predominant. R1

Br Br

2e–

R1

H

R2

+

Br

R1

Br

R2

H

(a)

R

2

(38.30)

(b)

R1 = H, R2 = Ph:

(a)/(b) = 1.6–5.3

R1 = Me, R2 = Ph:

(a)/(b) = 1.1–4.6

R1 = Me, R2 = COOH:

(a)/(b) = 1.9–6.1

R1 = Ph, R2 = COOH:

(a)/(b) = 1.0–10.1

R1 = Ph, R2 = COOEt:

(a)/(b) = 0.9–1.8

A similar diastereoselective reaction has been found in the reduction of 3-substituted 4,4-dichlorosuccinimides [106]. The stereochemistry of the products in an acidic medium is opposite to that in an alkaline medium. c. Bicyclic gem-Dihalides Bicyclic gem-dihalocyclopropanes always give predominantly the corresponding endo-monohalocyclopropanes, which seem to be formed by the reduction of the less hindered exo-halogen against the cathode in an adsorption state (Equation 38.31), although the isomeric ratio varies with the molecular structure of 6 and electrolytic conditions, such as solvent and supporting electrolyte [107,108]. H

H X

H

2e–

(38.31) X

X n

H 6

n

H 7

n = 2, X = Br: exo/endo = 0.2–0.9 n = 2, X = Cl: exo/endo = 0–0.5 n = 4, X = Br: exo/endo = 0

Similarly, cathodic reduction of bicylic gem-dibromocyclopropane (6) in the presence of chlorotrimethylsilane provides the exo-silylated isomer selectively (exo/endo=1.4–5.3) [109]. It is noticeable that an exo-bromo-endo-chlorocyclopropane (8a) normally gives the endo-chloride (9a) as a major isomeric product, but the endo-bromo-exo-chlorocyclopropane (8b) gives the exo-chloride (9b) and a mixture of 8a and 8b gives both isomers in an equal amount under the same conditions, as shown in Equation 38.32. Considering that alkyl bromides are more easily reduced than chlorides in general, this suggests that the reduction of gem-dihalocyclopropanes may occur at either halogen atom rather than only on the less hindered exo-halogen.

© 2016 by Taylor & Francis Group, LLC

1477

Stereochemistry of Organic Electrode Processes Cl

Cl 2e– H

Br

(38.32)

LiCl/EtOH exo-Br, endo–Cl (8a) endo-Br, exo–Cl (8b)

9a (endo/exo = 2.6) 9b (endo/exo = 0.38)

8a + 8b

9a + 9b (endo/exo = 0.97)

Similar stereochemical studies also have been made in detail using gem-dihalonorboranes and also bicyclic monohalides [110]. 2. Acyclic Polyhaloalkanes It was found that vic-dihaloalkanes were highly stereospecifically reduced at a mercury cathode to the corresponding olefins in high yields, as in Equation 38.33 [111]. The resulting stereochemistry suggests that the elimination of halogen atoms from the trans-periplanar conformation is strongly preferred. The detailed stereochemical course of the reaction was discussed by Brown and coworkers [112]. The stereospecific reduction of vic-dibromides to olefins was also reported [3,113]. H H H3C C C CH3

H

2e–

H3C

Br Br

H

(38.33)

CH3

dl, 90–95% (cis major or only) meso, 95% (trans only)

Monomeric and polymeric N-alkylpyridine-4-carboxylates serve as redox mediators for the indirect electroreduction of vic-dihalides to trans-olefins [114]. 1,1,2,2-Tetrachloroalkanes also give stereoselectively the corresponding 1, 2-dichloro-olefins in which the E-(trans)-isomers are major [115]. Trichloromethylmethanols are reduced with electrogenerated chromium (II) ion to give Z-monochlorovinyl compounds in good yields [116]. When a carboxy function is at the α-position, an E-double bond is predominantly formed. 3. Acyclic Halo-Olefins Monohalomaleic and fumaric acids and their esters are reductively dehalogenated with complete retention of the original molecular geometry in a wide range of pH, with the exception that bromomaleic acid gives mixtures of maleic and fumaric acids at pH 0.6–5.8 [117]. Indirect electroreductive cyclization of η-bromo-α,β-unsaturated esters using Co(III) or Ni(II) redox mediators provides six-membered ring products stereoselectively, as shown in Equation 38.34 [118]. EtO Br

O

EtO H COOMe

e– Ni(II) or Co(III) mediator

O H

MeOOC

E-isomer, 64–74% (cis/trans = 0.25) Z-isomer, 66–86% (cis/trans = 0)

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(38.34)

1478

Organic Electrochemistry

4. Intermolecular Dimerization a. Monohalides Ethyl α-chloro-phenylacetate was reductively dimerized and the d,l-meso ratio (0.4–1.4) of diethyl 2,3-diphenylsuccinate was affected by the cathode material [119] and a similar study was also reported by Rampazzo and coworkers [120,121]. They found that an excess of d,l-isomers derived from α-bromo-p-halogenophenylacetates decreases in the order of the F, Cl, and Br substituents in the phenyl group [120]. 3-Halo-6-nitro-5-cholestenes [122], aroyl chlorides [123], and benzylic halides [124] are also stereoselectively dimerized. b. Polyhalides Stereoselective dimerization of a gem-dihalide was achieved in direct and indirect electroreduction. α,α-Dibromotoluene was reductively dimerized by a Cr(II)/Cr(III) redox mediatory system to stilbene (cis/trans=0.18) [125], and direct electroreduction also gave similar results [124]. On the other hand, indirect cathodic dimerization of α,α-dichlorotoluene using a Co(II)(Salen) mediator gave a slight excess of cis-stilbene (cis/trans=1.1) [125]. α,α,α-Trichlorotoluene also gives 1,2-dichlorostilbene (cis/trans=1) [115]. A vic-dihaloadamantane undergoes the intermolecular reductive dimerization via unstable adamantene species to give stereoisomeric products in 80–90% yields, as in Equation 38.35 [126]. Cathodic reduction of α,α′-dibromo-1,2-dialkylbenzenes in the presence of dienophiles (cyclopentenedione, maleic anhydride, and N-phenylmaleimide, etc.) and hindered dienophiles provides Diels–Alder adducts stereoselectively [127,128]. In this reaction, the dienophiles play a dual role as mediators and dienophiles. X

H H

H

X 2e–

+

(38.35)

H 3:2

5. Intramolecular Cyclization The d,l- and meso-2,4-dibromopentanes are reductively cyclized, but neither the stereoselectivity nor stereospecificity is high [129]. On the other hand, 1,3-dibromo-1,3-dimethoxycarbonylpropane gives the corresponding trans-cyclopropane with high stereoselectivity as shown in Equation 38.36 [130]. Interestingly, the stereoselectivity was greatly affected by the solvent, cathode material, and applied potential. Br R dl or meso: Unknown: Unknown: Unknown:

Br

R

2e– R

R = Me: R = COOMe: R = COOMe: R = COOMe:

R

(38.36)

Hg/DMSO, 84–91% (cis/trans = 0.8–1.1) Pt/THF, 60–74% (cis/trans = 0.25) Pt/DMF, (cis/trans = 0.07) Hg/DMF, (cis/trans = 0.03)

The reduction of α,α′-dihaloketones may lead to the formation of stereoisomeric cyclopropanones, but their stereoselectivity could not be examined because of their instability under electrolytic conditions [131,132]. On the other hand, α,α′-dihalophosphinates were found to give stereoisomeric olefins via unstable oxyphosphirane species, but the stereoselectivity was low [133]. A similar result was also obtained in the reduction of α,α′-dibromodibenzylsulfone [134].

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1479

Stereochemistry of Organic Electrode Processes

I. OTHER HETEROATOM COMPOUNDS Only a few stereochemical studies on electroreduction of compounds containing heteroatoms other than oxygen, nitrogen, sulfur, and halogens have been reported so far [135,136]. Cathodically generated methyldiphenylgermyl anion having a tetrabutylammonium cation reacts with phenylacetylene to afford (Z)-2-methyl (diphenyl)germylstyrene predominantly. This is in sharp contrast to the same addition of methyldiphenyllithium giving the (E)-isomer preferentially [136].

J. ORGANOMETALLIC COMPOUNDS Stereoselective electroreduction of some organometal complexes has been reported. The reduction of α-oxotetramethylene-cyclopentadienyl-iron+-η6-benzenes in an acidic solution occurs specifically from the exo side to give only the endo-alcohol, as in Equation 38.37. In an alkaline solution, the corresponding stereoisomeric (d,l and meso) pinacols are formed, but the stereochemistry has not been determined [137].

R

O

Fe

2e–, +2H+

R

Fe

R

H OH R

(38.37)

R

R R = H, Me

endo-OH

IV. STEREOSELECTIVE AND STEREOSPECIFIC ELECTROOXIDATION A.

CARBOXYLIC ACIDS

1. Kolbe-Type Reaction a. Dimerization The stereochemistry of Kolbe dimerization products has been intensively studied [138,139], see Chapter 18. It is interesting to know if anodically generated radicals from carboxylic acids interact with the anode or not. It was found that the ratio of the isomeric dimers from ethyl phenylmalonate was not influenced by the kind of anode material [139]. Kolbe cross-dimerization has also been stereochemically studied mainly from a practical aspect with the aim of preparing stereoisomeric long-chain olefins from saturated and unsaturated carboxylic acids [140,141]. The stereoisomeric ratio (Z/E: 2.4–10.1) is influenced by current density and temperature. Although the pure Z acid was used as the starting material, the product was not the pure Z olefin but a mixture with the E olefin. This indicates that stereoisomerization around the double bond at the γ-position can occur in the Kolbe-type decarboxylation. b. Radical Addition to Olefins Unsaturated carboxylic acids can be decarboxylated to alkyl radicals that undergo an intramolecular addition. The 5-exo-trig-cyclization of β-allyloxy radicals, generated from an appropriate carboxylic acid, combined with a final cross-coupling has been applied to synthesize a precursor of prostaglandin PGF 2 α as shown in Equation 38.38 [142]. A radical tandem cyclization of a doubly unsaturated monocyclic carbocyclic acid provides a short synthetic route to tricyclic products, for example, triquinanes [143]. In all these reactions, the cyclization proceeds

© 2016 by Taylor & Francis Group, LLC

1480

Organic Electrochemistry

highly stereoselectively as intramolecular cis-addition, while the concluding intermolecular cross-coupling has a low stereoselectivity. CO2H AcO

O

OEt

R RCO2H –e

–, –CO

R

OEt 2

OEt +

O

AcO

O

AcO

(a)

(38.38)

(b)

R = CH3(CH2)7, 54%, (a),(b) = 3.0:1 R = PhMe2SiCH2, 38%, (a),(b) = 4.3:1

2. Non-Kolbe-Type Reaction a. Monocarboxylic Acids Carboxyl groups of cyclic monocarboxylic acids are anodically exchanged by nucleophiles, such as methoxide ion, to give stereoisomeric products [138,144,146]. 4-Substituted cyclohexanecarboxylic acids give the stereoisomeric methoxylated products, as in Equation 38.39 [138]. The reaction is nonstereospecific. Pyrrolidine-2-carboxylic acids are also methoxylated in high yields, but the stereoselectivity is generally low [144,145]. COOH R

OMe

–2e–, –CO2

(38.39)

R

MeOH

R = t-Bu, trans: 10.5–27.1% (cis/trans = 0.3–0.4) R = t-Bu, cis: 30.8% (cis/trans = 0.3) R = Ph, trans: 13.2% (cis/trans = 0.5)

Utley and Yates [147] reported an interesting stereochemistry and regiochemistry of methoxylated products derived from 4-phenylcyclohexen-2-carboxylic acid, as in Equation 38.40. Ph

COOH

–2e–, –CO2

Ph

+

Ph

OMe

(38.40)

MeOH MeO 26% (cis/trans = 0.4)

49% (cis/trans = 1.5)

Stereoisomeric (exo and endo) bicycloalkane carboxylic acids are stereospecifically methoxylated, hydroxylated, and amidated [147].

B. ACETOXYLATION Acetoxylation occurs in either of the following ways: (1) nucleophilic attack of acetate ion on cationic intermediates formed by loss of electrons or (2) radical attack of the acetoxy radical generated by loss of an electron from the acetate ion. 1. Olefins Diacetoxylation of monoolefins sometimes proceeds stereoselectively and/or stereospecifically [148,149]. The stereoselectivity is affected by the structure of the starting olefins and electrolytic conditions. 2. benzyl and Allyl Compounds The stereochemistry of acetoxylation at benzyl positions of substituted indanes and acenaphthenes was investigated in detail [150]. The cis–trans ratio of diacetoxylated products is significantly larger in the anodic processes than in chemical processes.

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It has been reported that the anodic acetoxylation at allylic positions proceeds with high stereoselectivity (88–93%), as does microsomal acetoxylation (88–93% stereoselectivity). A chemical method with perbenzoic acid gives selectivity as low as 74–80% [151]. 3. Cyclopolyenes Cyclooctatetraene gave stereoisomeric (cis and trans) bicyclo[4,2,0]octane-7,8-diols and their diacetates in water and acetic acid, respectively, and the anodic acetoxylation of anthracene resulted in the predominant formation of trans-9,10-diacetoxy-9,10-dihydroanthracene [152].

C.

METHOXYLATION AND HYDROXYLATION

1. Olefins Acyclic olefins are stereoselectively and sometimes stereospecifically dimethoxylated [153,154]. Anodic oxidation of 1,5- and 1,6-dienes in methanol provides the corresponding 1,4-dimethoxylated six- and seven-membered rings and the trans-isomers are favorably formed in both cases, as shown in Equation 38.41 [155]. Ph

Ph n

–2e–

MeO

MeOH

Ph

OMe

(38.41)

Ph n

n = 0, 64% (cis/trans = 0.04) n = 1, 52% (cis/trans = 0.04)

The anodic dimethoxylation of simple cyclic olefins, such as cyclohexenes, results in the predominant formation of trans-dimethoxycyclohexanes [154,156]. However, Barba and coworkers [157] reported that the stereochemistry of 1,2-dimethoxy-1,2-dihydroacenaphthene derived from acenaphthene is drastically influenced by the anode material. Palasz and Utley [158] have reported that N-acetylpiperidines are methoxylated via the enamide intermediate species formed by the oxidation. 2. Aromatics Aromatic hydrocarbons can also be anodically dimethoxylated to give the trans-dimethoxy dihydro derivatives in high or moderate stereoselectivity [159]. 3. Others Bicyclic cyclopropyl compounds are dimethoxylated to give stereoisomeric (cis/trans=1.1–1.3) cyclopropane ring-opened products [160]. Shono and coworkers found that anodic oxidation of borneol and isoborneol in methanol resulted in rearrangement to provide methoxylated stereoisomeric products with the same endo–exo ratio, as in Equation 38.42 [161].

OH

–2e–

O

MeOH endo or exo

(38.42) OMe

72–74% (endo/exo = 32)

Anodic oxidation of 4-substituted cyclohexanones in methanol in the presence of sodium halides in an undivided cell provided cis-5-substituted-2,2-dimethoxycyclohexanols in good yields as shown in Equation 38.43 [162]. OMe

–e– R

O

R NaBr or Nal/MeOH

R = Me, Et, t-Bu, Ph

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OMe OH

(38.43)

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Organic Electrochemistry

The α-positions of cyclic amines are stereoselectively methoxylated to provide cis-methoxylated product preferentially, as in Equation 38.44 [163]. Ar N

–2e–

COOEt

Ar N

MeO

COOEt

(38.44)

MeOH R

R R = trans-OH: 60%, cis/trans (MeO-COOEt) = 2.2 R = cis-OAc: 50%, cis/trans (MeO-COOEt) = 1.9

Noyori and Kurimoto [164] have found an interesting anodic ether exchange reaction of aryl glycosides. They described that this is a rare example of an SON1-type transformation reaction, but the reaction may be electrogenerated acid (EGA) catalyzed. The reaction is also stereoselective but not always stereospecific.

D.

FLUORINATION

Stereoselective anodic fluorination of various sulfur-containing cyclic compounds was studied by Fuchigami and coworkers [165–167]; see also Section V and Chapter 20. As shown in Equation 38.45, a fluorine atom was selectively introduced to the trans position regardless of the types of heterocycles. S

R X

O

–2e–, –H+

S

F

R

(38.45)

X

Et3N-nHF/MeCN

O 58–86%

X = NMe, R = 1-naphthyl: cis/trans = 29/71 (n = 3) X = O, R = p-CNC6H4: cis/trans = 39/61 (n = 4) X = S, R = mesityl: cis/trans = 20/80 (n = 4)

Interestingly ultrasonication changed the yield (also current efficiency) as well as stereoselectivity as shown in Equation 38.46 [168]. Without ultrasonication, the anodically generated cationic intermediate seems to be adsorbed on the anode surface, and fluorination takes place to result in the formation of the thermodynamically less favorable cis product considerably. However, under ultrasonication, the cationic intermediate should be desorbed from the anode surface, followed by fluorination to provide the thermodynamically favorable trans product preferentially. S

Ph N

O

Me

–2e–, –H+ Et3N-3HF 3F

S

F O

N Me trans

S

F

Ph +

O

Ph N Me

(38.46)

cis

Under mechanical stirring, 24% (cis/trans = 50:50) Ultrasonication, 77% (cis/trans = 38:62)

Intramolecular carbon–carbon bond formation followed by fluorination proceeds as shown in Equation 38.47 [169]. It is notable that the cis forms are always major, and the cis–trans ratio depends on the nature of the electroauxiliaries (EA).

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1483

Stereochemistry of Organic Electrode Processes EA –ne–

O

O

Bu4NBF4

C7H15

O

(38.47)

+ C7H15

C7H15

F cis

F trans

EA = SiMe3, 68% (cis/trans = 55/45) SnBu3, 98% (cis/trans = 74/26) SMe, 64% (cis/trans = 87/13)

E. ACETAMIDATION The anodic acetamidation of olefins proceeds smoothly in wet acetonitrile containing suitable nucleophiles leading to stereoisomeric products. 1. Haloacetamidation The stereoselectivity of the fluoroacetamidation of a variety of olefins has been rationalized as due to adsorption of the olefin, fluoride anion, and the resulting oxidized olefinic species on the anode surface as shown in Equation 38.48 [170–172]. Ph

R2

–2e–

R1

R3

F–/MeCN

Ph R2 AcHN

C C F R1

(38.48)

R3

R1 = R2 = H, R3 = Ph (erythro/threo = 0.28) R1 = R3 = H, R2 = Ph (erythro/threo = 0.25) R1 = R3 = Me, R2 = H (erythro/threo = 3.00)

Chloroacetamidation of unsaturated carbohydrates occurs with substrate-dependent stereoselectivity [173]. For iodoacetamidation of cyclohexene [174], a trans-selectivity has been reported. 2. Thioacetamidation Bewick and coworkers [175] found that the anodic thioacetamidation of cycloalkenes occurred highly stereoselectively in acetonitrile containing disulfides, as in Equation 38.49. NHAc

–2e–

(38.49) n

RSSR/MeCN

n

SR

n = 1, R = Ph, 56% n = 2, R = Ph, 50% n = 2, R = Me, 58%

F. OTHER ADDITION REACTIONS TO OLEFINS 1. Halogenation Olefins can be anodically dihalogenated, sometimes with very high stereoselectivity. a. Difluorination Laurent and coworkers [171] reported the difluorination of 3-t-butylindene, but the stereoselectivity was not high. On the other hand, Fuchigami and coworkers found that anodic difluorination of 3-substituted benzofuran derivatives and 1-acetyl-3-substituted indole derivatives provided cis- and transdifluorinated products as the main stereoisomer, respectively, as shown in Equation 38.50 [176,177].

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1484

Organic Electrochemistry CN

F

CN

–, +2F–

F

Z=O

O

F –2e–, +2F–

–2e

F

Z = NAc

Z

CN

N Ac trans only

cis/trans = 2.2

(38.50)

b. Dichlorination and Dibromination The dichlorination of olefins, dienes, and acetylenes is highly stereoselective and gives almost exclusively single stereoisomers of products, as in Equations 38.51 and 38.52 [178]. The bromination also gives single stereoisomers of vic-dibromo compounds [179,180]. Ph H

H

Cl

Cl

C

C

H

H

–2e– Ph

Cl–

Ph

(38.51)

Ph

meso 65%

R1

R2

–2e–

R1

Cl

Cl–

Cl

R2

(38.52)

trans R1 = CH2OH, CH(OH)Ph, R2 = H, 60–70% R1 = R2 = Ph, 40%

2. Selenoalkoxylation, Thioalkoxylation, Thiohydroxylation, and Disulfenylation The selenoalkoxylation of olefins, such as cyclohexene, cyclopentadiene, and indene, proceeds highly stereoselectively and regioselectively to give trans products in high yields [181]. Recently, stereoselective thiohydroxylation, thioalkoxylation, and disulfenylation of cyclic olefins were achieved by low temperature anodic oxidation of ArSSAr generating ArS(ArSSAr)+ followed by reaction with 1-methylcyclohexene. In all cases, anti addition products were obtained exclusively since these products were formed via episulfonium ion intermediate as shown in Equation 38.53 [182].

OMe MeOH

68% SAr

–e–

H2O

[ArS(ArSSAr)+BF4–]

ArSSAr Bu4NBF4/CH2Cl2 –78°C

OH 79%

–78°C

(38.53)

SAr SiMe3 SAr 58% SAr

G. INTERMOLECULAR AND INTRAMOLECULAR REACTIONS OF OLEFINS AND/OR AROMATICS 1. Dimerization Olefins [183] and aromatics [184] are oxidatively dimerized to give stereoselectively the dienes and biaryls (rotamers), respectively.

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Stereochemistry of Organic Electrode Processes

Maleic anhydride is stereoregularly polymerized by anodic initiation to give mainly cis-poly (vinyleneketoanhydride) [185]. 2. Cross-Coupling The intramolecular coupling of enol ethers with enol ethers, styrenes, dienes, alkyl-substituted olefins, allylsilanes, and vinylsilanes was systematically studied by Moeller [186,187]. These couplings allow the smooth formation of quaternary carbon atoms in fused bicyclic rings having a cis stereochemistry (Equation 38.54) [188,189]. MeO –2e– MeOH/THF lutidine

TMS

TBDMS O

O TBDMS

MeO

(38.54)

MeO

75%

Ketene dithioacetal enol ethers underwent stereoselective intramolecular coupling to yield transor cis-disubstituted five- or six-membered cycloalkanes with high stereoselectivity and yield [189]. In the intramolecular coupling of an allylsilane, a five-membered ring with three contiguous stereogenic centers was stereoselectively formed without loss of a very acid-sensitive allylic alkoxy group [190]. The intramolecular coupling of a furan ring with a silylenolether was used to stereoselectively assemble the tricyclic core ring system of alliacol A as shown in Equation 38.55 [189]. TBS

O

O –e–

TBS

(38.55)

MeOH/CH2Cl2 lutidine

OO

H

TsOH O HO 88%

Anodically generated phenoxy cations, o-quinones, and o-quinone methides react with olefins to give bicyclic and tricyclic annelated compounds stereoselectively [191–194]. Anodic induced intermolecular [2 + 2] cycloaddition reactions between an aliphatic cyclic enol ether and unactivated olefins have been demonstrated to proceed diastereoselectively as in Equation 38.56 [195,196]. PhO

+1.0 V vs. Ag/AgCl

+ O

1 M LiClO4/MeNO2 0.5 F

PhO

(38.56) O cis only

H. ACTIVE HYDROGEN COMPOUNDS (CARBON ACIDS) Only few studies of the stereochemistry of anodic oxidative dimerization of active hydrogen compounds (carbon acids) and their salts (carbanions) have been reported. The anion of methyl phenyl acetate, formed by an electrogenerated base, was homocoupled with iodine or anodically mediated by iodide to afford dimethyl 2,3-diphenylsuccinate in high yield and high d,l-selectivity. This reaction probably does not involve free radicals but an iodination-nucleophilic substitution sequence [197,198].

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Organic Electrochemistry

Electrolysis of malonate and alkylidenecyanoacetates in alcohols in the presence of NaBr resulted in the stereoselective formation of (E)-3-substituted 2-cyanocyclopropane-1,1,2-tricarboxylates in good yields (Equation 38.57) [199]. The reaction involves both anodic oxidation of Br− and cathodic reduction of malonate generating the corresponding anion. R1

CN

(COOR2)

CH2

+

R1 H

Electrolysis 2

R2OOC

NaBr/R2OH

COOR2

CN COOR2

R2OOC

(38.57)

R1 = Me, Et, Pr, Ph R2 = Me, Et

I. HETEROATOM COMPOUNDS 1. Nitrogen Compounds A few interesting reactions of nitrogen compounds from a stereochemical aspect have been reported. Tokuda and coworkers [200] reported a stereochemically clear result that lithium alkenylamides undergo a highly stereoselective cyclization by anodic oxidation to give cis-1methyl-2,5-disubstituted pyrrolidines exclusively in reasonable yields, as in Equation 38.58. The cyclization seems to proceed via neutral amino radicals adsorbed on and/or near the anode surface. –2e–

R –

N Me

R

Li+

N

(38.58)

Me

Me cis R = Me, Et, 4-MeC6H4, 31–52%

2. Sulfur Compounds Anodic oxidation of saturated cyclic sulfur compounds has been studied from stereochemical aspects. Kimura and coworkers [201] reported that a 4-substituted thiane was stereoselectively oxidized to the trans-thiane oxide, which was always the major isomeric product, although the isomeric ratio was affected by electrolyte and solvent used, as in Equation 38.59. The oxidation of 2,5-substituted-1,3-dithianes to the cis- and trans-4-substituted-1,2-dithiolanes is stereoselective but not stereospecific [202]. S

p-ClC6H4

O

–2e–

S

(38.59)

p-ClC6H4 50% (cis/trans, 0.08–0.52)

Anodic oxidation of ketene dithioacetal having an aminoalkyl group in methanol provided the cyclized product stereoselectively as shown in Equation 38.60 [203]. S

S Me H2N

S

–e

–

Me

Me + H2N

S

–H+

H N

S S Me

Me Me

© 2016 by Taylor & Francis Group, LLC

–e– MeOH

HS N

S OMe Me

Me

(38.60)

Stereochemistry of Organic Electrode Processes

1487

3. Silicon Compounds Yoshida and coworkers [204] found that stereoisomeric mixtures of 1-trimethylsily-1,3-dienes are oxidized in methanol to give trimethoxylated trans-olefins as single stereoisomers. Anodic oxidation of 1,2-bis(trimethylsilyl)-1,2-diphenylethane in methanol afforded 1,2-dimethoxy-1,2-diphenylethane (d,l/meso=10) [205]. γ-Acyloxyvinylstannanes undergo cupper-mediated anodic homocoupling to afford 1,3-dienes without isomerization. Thus, optically active 1,3-dienes were prepared anodically from optically active alkenylstannanes [206].

J. ORGANOMETALLIC COMPOUNDS Stereoselective synthesis of organometallic complexes has been achieved in the oxidative addition of aryl halides to triethylphosphine nickel(0) complexes, leading to the exclusive formation of trans-aryl nickel(II) halide complexes [207]. Electron-transfer reactions of the Fe of cis- and trans[η-C5H5Fe(CO)SR]2 occur stereospecifically with no stereoisomerization on changing the oxidation state of the Fe [208]. In the electrochlorination of a ligand (R) of the η-C5H5Fe(CO)R complex, the stereochemistry is retained [209].

V. ASyMMETRIC REDUCTION AND OXIDATION A. REACTION OF CHIRAL COMPOUNDS Electrochemical reactions that occur at chiral centers and also give chiral products are mentioned here. 1. Reduction of Halides Cathodic reduction of chiral (optically active) haloalkanes to the chiral alkanes has been studied mainly from a mechanistic aspect. The configurational change (retention and inversion) during the reduction is discussed. The reduction of chiral 1-substituted 1-halo-2,2-diphenylcyclopropanes at a mercury cathode to the chiral cyclopropanes has been reported by some workers [210,211]. The 2e reduction of (S)-(+)-1-bromo-1-carboxy-2,2-diphenylcyclopropane showed that the stereoselectivity at an Hg cathode was strongly determined by the supporting electrolyte cation. With NH4+ a preferential retention of configuration was observed, which increased with a more negative reduction potential. In contrast, an R4N+ cation gives rise to a major inversion, which increases with the bulkiness of the organic cation. The stereochemistry is thought to be determined by the facial selective protonation of the final anion, whose orientation and shielding is also controlled by the nature of the interface [211]. The reduction of chiral open-chain halides has been also reported. Nonaka and coworkers [212] investigated the reduction of a chiral chloroalkane and attempted to rationalize the stereochemical course of the reaction by considering the configurational stability of the radical intermediate species at the cathode interface. Optically pure (+)-(1R,2R)-1,2-dimethylcyclopropane was obtained in the reductive intramolecular cyclization of (+)-(2S,4S)-2,4-dibromopentane. Maran [213] investigated electroreduction of optically active 2-bromo-N-phenylpropanamide giving optically active N-phenyl-2-hydroxypropanamide and cis-1,4-diphenyl-3,6-dimethyl-2, 5-dioxopiperazine. 2. Reduction of Alcohols and Amines Electroreduction of alcohols and amines hardly occurs in general. Only peculiar alcohols and amines can be reduced to the corresponding alkanes.

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Optically active O-benzoylatrolactic acid and its methyl ester were reduced to 2-phenylpropionic acids with almost complete loss of optical activity. On the other hand, Nonaka and coworkers [214] found that 1-(2-pyridyl)alkanols were reduced to 2-alkylpyridines with configurational retention in an acidic medium, while 1-(4-pyridyl)alkanols gave racemic 4-alkylpyridines. The stereochemical course was discussed considering perpendicular adsorption of pyridylalkanol molecules on the cathode and the configurational stability of carbanion intermediates. Electroreductive deprotection of tosylates and sulfonamides of chiral alcohols and amines, respectively, was achieved without racemization. 3. Reduction of Onium Compounds Optically active quaternary arsonium [215,216] and phosphonium salts [217] are cathodically cleaved to tertiary arsines and phosphines, respectively, with retention of configuration. The first enantiomer enriched chiral phosphines have been prepared using this method. 4. Oxidation of Carboxylic Acids and Alcohols Optically active carboxylic acids with achiral center at the α-carbon provide racemic Kolbe dimers [218,219]. Racemization is also observed in anodic methoxylation of carboxylic acids [220]. These facts may suggest that Kolbe intermediate radicals interact so weakly with the anode surface that their configuration cannot be stabilized during the reaction. Matsumura, Onomura, and coworkers developed unique memory of chirality [221–223]. They found that when optically active N-o-phenylbenzoylated oxazoline and thiazoline derivatives were anodically oxidized, optically active methoxylated products were formed with 83% and 91% enantiomeric excess (ee), respectively, as shown in Equation 38.61 [221,222]. O Me Me

N O

Me Me

O Me Me

COOH or

N

O

COOH –2e

O

–

NaOMe –30°C

Me Me

O Me Me

OMe

N

or

O

Me Me

N

OMe

(38.61)

O

83% ee

91% ee

The memory of chirality was similarly observed in the anodic methoxylation of optically active amino alcohol derivatives, as in Equation 38.62 [223]. Ar Ar OH

N O Ph

–2e– NaOMe –30°C

OMe

N

Ar

Ar O

+

(38.62)

O Ph 73% ee

β-Chiral carboxylic acids give optically active dimers in high optical purity [224,225]. 5. Oxidation of Heteroatom Compounds Shono and coworkers [226] reported that an α-chiral alcohol was converted into the corresponding chloride by diphenylsulfide-mediated anodic reaction. Anodic cyclization of α-amino acids in methanol was also reported [227]. Ohtani and coworkers [228–230] found photocatalytic one-step synthesis of l-proline and l-pipecolinic acid from l-ornithine and l-lysine using aqueous semiconductor suspensions.

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B. ASYMMETRIC SYNTHESIS Asymmetric synthesis is a highly important area in organic chemistry. However, it must be noted that electrochemical asymmetric synthesis has been still immature compared with well-established chemical, catalytic, and enzymatic methods, although much effort has been made since the first report in this area by Grimshaw and coworkers [231] in 1967. 1. Use of Chiral Electrode Adsorbants In most studies, optically active alkaloids have been used as asymmetry inducers in this method. The optical yield was found to be greatly dependent on both molecular structures of substrates and alkaloids as follows. Asymmetric reduction of activated olefins was performed using prochiral 4-methylcumarin [231–233] and a maximum asymmetric yield reported was 17% as shown in Equation 38.63 [232]. Later, Schäfer and coworkers carried out similar enantioselective electroreduction of various 4-substituted coumarins by systematic variation of the electrolysis conditions, and they obtained optical yields as high as 67%, Equation 38.63 [234–236]. O

O

2e–, +2H+

O

O

(38.63)

Hg cathode H Me

Me

alkaloid = sparteine (0.07 equiv.), pH = 5.4, 17% ee alkaloid = yohimbine (0.01–0.05 equiv.), pH = 2–3, 47–67% ee

The asymmetric reduction of prochiral ketones like acetophenone and various acetylpyridines to the corresponding chiral alcohols has been extensively investigated [237–245], and 48.4% ee has been achieved from 2-acetylpyridine using strychnine salt as a supporting salt [238]. Later, Yadav and coworkers carried out asymmetric reduction of acetophenone in aqueous DMF containing N,Ndimethylephedrinium tetrafluoroborate and ethynyl pentyl ketone in i-PrOH/DMF containing DMQ tetrafluoroborate to give 55% ee and 70% ee, respectively, as shown in Equation 38.64 [246,247]. 2e–

O (CH2)4CH3

DMQBF4 Bu4NBF4 DMF/i-PrOH

OH

(38.64) (CH2)4CH3

72% 70% ee

The asymmetric reduction of C=N double bonds has been investigated using prochiral oximes [244,245,248–250]. The maximum enantiomeric excess is 18% [245]. gem-Dihalides are also asymmetrically reduced to chiral monohalides [251,252], and a maximum enantiomeric excess of 44.3% has been obtained in the cathodic reduction of 1,1-dibromo-2,2diphenylcyclopropane in the presence of emetine (5 × 10 −4 M) [252]. Not only asymmetric yields of products but also their absolute configurations vary drastically with cathode potential, supporting electrolyte, solvent, pH, cathode material, concentration, and other factors that affect adsorption [253–256]. It may sometimes be a disadvantage in asymmetric synthesis that a reaction is too sensitive to a variety of conditions. 2. Use of Chiral Supporting Electrolytes Although a large amount of chiral compounds are required, this method was applied to the asymmetric reduction of ketones [257–259] and imines [258]. Maximum asymmetric yields reported for the former and the latter are 20% [257] and 8.95% [258], respectively. A higher asymmetric yield

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Organic Electrochemistry

(20.6%) was obtained in the hydrodimerization of a ketone [259]. It is a problem that lower asymmetric yields are obtained using much larger amounts of asymmetry inducers, which must play the role of supporting electrolyte. Maekawa, Nishiguchi, and coworkers achieved asymmetric anodic acetoxylation (44% ee) of enol ester derivative using camphorsulfonic acid salt as a chiral supporting electrolyte [260]. 3. Use of Chiral Media Seebach and Oei [261,262] reported the asymmetric hydrodimerization of acetophenone (a maximum asymmetric yield of 6.4%) in a chiral cosolvent. It is interesting that in the presence of β-cyclodextrin, head-to-tail coupling of acetophenone leads to optically active dimeric monoalcohol (ca 24% ee), whereas the head-to-head coupling (pinacol formation) gives optically inactive pinacol [263,264]. 4. Use of Chiral Oxidant Electrochemical Os-catalyzed asymmetric dihydroxylation of olefins was achieved using Sharpless’ ligand [265–267]. 5. Use of Chiral Modified Electrodes The use of chiral modified electrodes may be the most promising method for electrochemical asymmetric synthesis because of (1) use of extremely small amounts of inducers, (2) low oversensitivity to reaction conditions, (3) variety of kind, and (4) possible application to both cathodic and anodic reactions, although there are some problems to be solved. a. Chemically Modified Electrodes Miller and coworkers [268,269] found that prochiral carbonyl compounds were reduced to chiral alcohols on a cathode chemically modified with (S)-phenylglycine. This may be the first example of a chemically modified electrode (CME). The maximum enantiomeric excess was 14.5% in the reduction of 4-acetylpyridine. A similar electrode modified with (+)-camphoric acid was used for the oxidation of p-tolyl methylsulfide to the sulfoxide with 2.5% ee [270]. However, no asymmetry induction could be observed in the reduction of carbonyl compounds on other modified cathodes with α-chiral amines [271]. Osaka and coworkers studied intensively chiral discrimination with electrodes modified with an enantiomeric amino acid as follows [272–276]. They achieved enantioselective crystal growth of leucine on a self-assembled monolayer (SAM) with covalently attached leucine molecules [272]. They also demonstrated that the gold electrode modified with SAM of homocysteine exhibited enantioselectivity in the redox reaction of 3,4-dihydroxyphenylalanine (DOPA) in acidic solution [273]. Furthermore, an extremely enhanced enantioselectivity was achieved for the detection of enantiomers of alanine, leucine, and DOPA based on the voltammograms for the deposition of Cu from Cu complexes of the amino acids at an Au electrode modified with SAM of l-homocysteine [274,275]. In addition, the stereospecificity in redox reactions of molecules with two chiral centers, catechin and epicatechin, at the gold electrode modified with the homocysteine SAM was investigated, and they found that the homocysteine SAM exhibits a chiral recognition ability to one of the two chiral centers. These results are useful for the design of electrodes and catalysts for chiral sensing as well as asymmetric synthesis [276]. b. Electrodes with Irreversibly Adsorbed Inducers Osa and coworkers [277] developed a chiral Raney nickel electrode on which (S,S)-(−)-tartaric acid was irreversibly absorbed and applied it to the asymmetric reduction of carbonyl compounds. A maximum asymmetric yield of about 20% was obtained in the reduction of methyl acetoacetate.

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1491

Stereochemistry of Organic Electrode Processes

c. Polymer-Coated Electrodes Nonaka and coworkers [278–283] prepared a variety of electrodes coated with optically active poly(amino acid)s and applied them to both asymmetric reduction and oxidation. Maximum enantiomeric excess values in the reduction of 4-methylcumarin and oxidation of sulfides were 43% [278] and 93% [283], respectively. However, later, the reproducibility of ee value in the sulfide oxidation was found to be greatly affected by prepared poly(amino acid)s. Osa and coworkers [284,285] developed a graphite felt electrode modified with 2,2,6,6-tetramethylpiperidin-1-yloxyl (TEMPO) and applied it to enantioselective, electrocatalytic oxidative coupling of naphthol, naphthyl ether, and phenanthrol in the presence of (−)-sparteine as a base. The enantioselectivity of the coupling products was very high as 98%. d. Clay-Metal Chelate-Coated Electrodes Yamagishi and Aramata [286] prepared a clay-coated electrode that incorporated optically active tris(1,10-phenanthroline)ruthenium (II) into the montmorillonite thin film. In the oxidation of sulfides, 20% ee was obtained. e. Intercalated Graphite Electrodes Simonet and coworkers [287] prepared a new type of chiral electrode by forming a lamellar compound (CnNR4*) of graphite and optically active ammonium by electrochemical intercalation. This electrode provided 9.3% ee as a maximum value in the reduction of phenyl t-butyl ketone to the alcohol. 6. Intramolecular Asymmetric Induction Intramolecular asymmetry induction is not a unique method characterized by electrochemistry. However, a number of papers dealing with this methodology have been reported to date. The reduction of optically active alcohol esters or amides of α-keto acids [288,289], oximes [290], and the oxidation of olefins [291] have been reported. Schäfer and coworkers achieved 81% ee in the reduction of (S)-4-isopropyl-2-oxazolidinone phenylglyoxylate [289]. Waldvogel and coworkers developed an efficient method for electrosynthesis of menthylamine from menthone oxime [290]. The reduction of menthone oxime at a mercury cathode at −10°C provided (−)menthylamine selectively in 76% yield (d.r.=4.1:1) (Equation 38.65). When a lead cathode was used, its corrosion took place. However, addition of a small amount of quaternary di-, tri-, or tetra-ammonium salt to 2% H2SO4/MeOH provided (+)-neomenthylamine selectively in quantitative yield (d.r.=0.5–0.6:1) without corrosion of the lead cathode. Even though the product is racemic, pure menthylamine can be separated after the diastereomeric mixture (1:1) is converted to the hydrochloride salt followed by separation owing to the different solubility in t-BuOMe. This new method is highly useful [290].

4e–, +4H+

+ NH2

NOH

(a)

NH2

(38.65)

(b)

In 2% H2SO4/10% aq. DME, –10°C, Hg cathode, 76% (a:b = 4.1:1) In 2% H2SO4/MeOH, 20°C, Pb cathode, 99% (a:b = 1:2)

Nonaka and coworkers [292] found that amino acid N-carboxy anhydrides were polymerized with various electrogenerated bases as catalysts to give the poly(amino acids) with highly chiral secondary structures in high yields. Conducting chiral poly(thiophenes) prepared by electropolymerization were used for chiral anion recognition[293].

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1492

Organic Electrochemistry

It is known that 8-methylene-oleandomycin is reduced by catalytic hydrogenation (H2, Pd/C) and provides a mixture of antibacterially active (8R)-methyl-oleandomycine and antibacterially inactive (8S)-form epimers, whereas catalytic hydrogenation using Raney/Ni results in the (8R)-methyl epimer as a major product. In sharp contrast, cathodic reduction of 8-methylene-oleandomycin at an Hg cathode gave the (8S)-product as a main product (d.r.=5:1) (Equation 38.66) [294]. The stereoselectivity can be explained by the bottom-side adsorption of the substrate through its conjugated double bond to the cathode surface.

OH

HO O O

O O

(8S)

O

O

O

O

OH

2e–, +2H+

N

OH

HO O O

O O

O

N O

OH

(8R)

O

HO O O

O O

O

O

OH

O

O

N

(38.66)

OH

O

Inactive

Antibacterially active

Moeller and coworkers [295] found that electrochemical cyclization (anodic intramolecular alkoxylation) of optically active dipeptides proceeded highly diastereoselectively. Steckhan and coworkers [296,297] reported that anodic methoxylation of chiral 5-methyl- and 5-chloromethyl-2-oxazolidinones followed by Lewis acid catalyzed allylation provides 4-allyl products highly diastereoselectively. Similarly, anodic methoxylation of cyclic dipeptides and dipeptolides derived from chiral α-amino acids [298] or α-hydroxy acids [299] provides useful chiral synthetic building blocks, as in Equation 38.67.

N H

N

* COOH

*

O

–2e–, –CO2, +H+

NH

O

N

MeOH

O

COOH

SiMe3

* NH OMe

O

N O

O

*

Hydrolysis NH

N H

* COOH

+

*

COOH

(38.67)

NH2

99% de

It is interesting that anodic methoxylation of optically active N-protected α-amino acids with one more chiral center and subsequent chemical alkylation affords the final products in high optical purity, although the diastereomeric purity of the intermediate methoxylated products is not so high [300]. In these cases, intramolecular asymmetric induction may occur in the chemical alkylation step. Similarly, highly diastereoselective allylation of phosphonamide derivatives having phosphoryl chiral auxiliary was achieved as shown in Equation 38.68 [301].

Me

O N P O O Et Ar

–2e–, +H+

Ar = 2,6-Cl2C6H4

© 2016 by Taylor & Francis Group, LLC

MeOH

OMe O Me N P O Et O Ar dr = 65/35

SiMe3 Me Bu2BOTf –78°C

O N P O O Et Ar dr = 94/6

(38.68)

1493

Stereochemistry of Organic Electrode Processes

Anodic oxidation of bicyclic amine prepared from (S)-prolinol and trifluoroacetaldehyde provided an enantiomerically pure methoxylated product in excellent regioselectivity. The methoxylated product was readily transformed into (S)-α-allylprolinol (Equation 38.69) [302]. OH OEt

F3C N OH H

MeO

–2e– N

p-TsOH

CF3COOH

N MeOH

O

O

+

N O

CF3

CF3

SiMe3

CF3

HCl

N

MeOH

O

OH

CF3

(38.69)

N H

Shankaraiah and coworkers anodically generated the pooled N-acyliminium cation from cyclic N-carbamates having an 8-phenylmenthyl group at 0°C, followed by addition of n-Bu3SnCN this provided the α-cyanated product in 40% yield with 85% diastereomeric excess (de) [303]. When TMSCN was used instead of n-Bu3SnCN, the de decreased to 65%. Interestingly, addition of catalytic amounts of β-cyclodextrin markedly increased the de up to 91% as shown in Equation 38.70 (see also Chapter 37). CN O

–2e–

+

, –H

N OR*

+

O

CN

N

low T

O

–

OR*

R = 8-phenylmenthyl

(38.70)

N OR*

“Cation pool”

Onomura and coworkers also achieved highly diastereoselective anodic cyanation of an l-proline derivative, as in Equation 38.71 [304]. –2e–, +CN– MeO2C

N Tr Tr = triphenylmethyl

MeO2C

N

CN

(38.71)

Tr 92% (cis only)

3,3,3-Trifluoropropylsulfides having a methoxy or an ester group at the 2-position undergo anodic α-methoxylation efficiently in the presence of fluoride ions. In these cases, the diastereoselectivity of the products is moderate [305,306]. Kise and coworkers [307] reported electroreductive intramolecular hydrocoupling of diimines derived from chiral 1,2-diamines giving enantiomerically pure piperazines. They also prepared enantioselectively 3,4-diphenyl adipate by electroreductive intermolecular hydrocoupling of chiral N-trans-cinnamoyl-2-oxazolidones [308]. Carboxylates with different nonracemic chiral auxiliaries have been anodically decarboxylated to explore the face-selective heterocoupling of the intermediate radicals. With (2R,5R)-2,5dimethylpyrrolidine as the chiral amido group and an increasing size of substituents R1 and R2 selectivities, up to 86% de have been obtained as shown in Equation 38.72 [309,310]. The results point to an increasing portion of an intermediate radical with Z-conformation and a growing steric hindrance for the si-approach. With oxazolidine, 2,10-camphersultame, and menthol auxiliaries, the diastereoselectivity is lower. Cross-coupling of 2-carboxy-butyrolactones with a stereogenic center at the α-position to the intermediate radical leads to 2-substitued butyrolactones in 33–43% yield and up to 88% de, as in Equation 38.73 [311].

© 2016 by Taylor & Francis Group, LLC

1494

Organic Electrochemistry O

O N

O OH +

–e–, –CO2

R2COOH

N

1

+ R2 R1

R

R1 = CH2C6H5, C(CH3)3, C(CH3)2 Et R2 = C4H9, CH2C(CH3)3, C(CH3)2 Et O

O R2

N

R2

N

+

(38.72)

R1

R1 16–69% (20–86% de)

R1 R2 R2

R

2H

R2 R2

–H+, –e–, –CO2

O

O

3CO

R1 = Ph, i-Pr, t-Bu R2 = Me, H R3 = CH3, C(CH3)2CO2Et

O

R3

R1

R3

R1

CO2H

R2 R2

+

O

O

O

(38.73)

cis major trans major 33–43% (0–88% de)

The results show that in coupling of anodically generated radicals, good diastereoselectivities can only be achieved when the intermediate radical adopts a preferred conformation and the substituents efficiently shield one face of the radical. These observations correspond to those found in homogeneous radical reactions [312]. Stereogenic centers at the β-carbon atom of the acid and more remote positions retain their configuration as expected. The same holds for double bonds except for (Z)-4-enoic acids that partially isomerize to (E)-products via a reversible ring closure to cyclopropylcarbinyl radicals [140]. Schäfer and coworkers [313] reported interesting diastereoselective cross-coupling of anodically generated radicals bearing chiral amide groups. A highly asymmetric induction in the nickel-catalyzed electroreductive coupling of aryl halides with α-chloropropionic acid derivatives was also reported by Nedelec and coworkers, as sown in Equation 38.74 [314]. O ArBr +

Cl

O

2e– OR*

Me

Ar

NiBr2, bipy Al anode

(S) OR*

(38.74)

Me

O R* =

N (S) Ph

Me N– (R) Me

85% ee (93% de)

Ar = MeO

Ar =

Br

82% ee (92% de) F

Asymmetric synthesis of α-fluorinated cyclopropylphosphonamides was successfully carried out by the cathodic reduction of chiral dibromofluoromethylphosphonamides in the presence of t-butyl acrylate, as in Equation 38.75 [315].

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1495

Stereochemistry of Organic Electrode Processes

— —

N

O

R1 P

Mg anode

CFBr2

–

R1

H

R1

(38.75)

trans (40% de) R1

CO2R2

N

R1

— —

N

H

R1

2e–

CO2R2

N

O

N

F

P

P

H

O F H

N

N

–

–

R1 = CH2C(CH3)3, R2 = t-Bu

— —

O

F

–

— —

N

O

P

R2

CO2

–

N

N

F

P

— —

R1

R1

CO2R2

R1

R1

CO2R2

cis (95% de) cis/trans = 1/9

Feroci, Inesi, and coworkers studied diastereoselective electrochemical carboxylation of α-bromocarboxylic acid derivatives having various chiral auxiliaries. When Oppolzer’s camphorsultam was used as a chiral auxiliary, good yield (80%) and excellent diastereoselectivity (98:2) were obtained, as in Equation 38.76 [316].

e–, +CO2

O S O

O

N

CH2N2

S O

Br

O

O

+

N

O

S O

CO2Me

(38.76)

N CO2Me

O

83% d.r. =2:98

They also carried out similar diastereoselective cathodic carboxylation using cinnamic acid derivatives having various chiral auxiliaries as shown in Equation 38.77 [317]. When 4-R-(diphenylmethyl)oxazolin-2-one was used as a chiral auxiliary, the R-isomer was formed as a major product (d.r.=7:3), and the two diastereomers were easily separated by flash chromatography. O N

Ph Ph

Ph

O

MeOOC

O O

2e–, +CO2 CH2N2

Ph (R)

O N

O

(38.77)

Ph Ph Major product

Synthesis of chiral compounds having a fluorine at a chiral center has been becoming of much importance in medicinal chemistry and material science. However, asymmetric electrochemical fluorination is quite difficult in general, due to the small size of a fluoride ion and the necessity of use of polar solvents for electric conductivity [318]. Fuchigami and coworkers studied diastereoselective anodic fluorination of α-(phenylsulfenyl)acetic acid esters having various chiral auxiliaries as an ester moiety [319]. However, even when the 8-phenylmenthyl group was used as the chiral auxiliary, the diastereoselectivity was low (~28% de). The low diastereoselectivity is due to the long distance between the reaction site and the chiral auxiliary moiety. Then, they investigated diastereoselective

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1496

Organic Electrochemistry

anodic fluorination of sulfides having two hydroxy groups at β- and γ-positions via 1,2-asymmetric induction as follows [320]. As shown in Equation 38.78, the diastereoselectivity was greatly affected by protecting groups of the diol moiety, and the 2-spirocyclohexyl-1,3-dioxolan-4-yl group gave the highest diastereoselectivity. The 1,3-dioxolanyl and 2,2-dimethyl-1,3-dioxolanyl groups gave ca. 70% de, while the 1,3-dioxolanonyl group decreased the de drastically to 13%. Similarly to the case of the 1,3-dioxolanonyl group, the oxazolidinone group was not effective for diastereoselective fluorination of sulfide [321].

O

R=

R

R

–2e–, –H+

O

O

Et3N-3HF/MeCN 2F

SPh

54% (80% de)

O SPh

(38.78)

Me Me

R=

45% (70% de)

F H H

R=

51% (68% de)

(1R,2S) isomer: major product O R=

22% (13% de)

Highly diastereoselective anodic fluorination of chiral 1,3-oxathiolanones derived from camphorsulfonamides was achieved as shown in Equation 38.79 [322]. A single diastereomer was always obtained regardless of bulkiness of the N-alkyl groups on the sulfonamide and electrolytic conditions. The absolute configuration of starting materials is (1S,2R), while that of the fluorinated carbon of the products is the S-form. Therefore, it is reasonable that a fluoride ion predominantly attacks the anodically generated cationic intermediate from the less hindered re face, because the sulfonamide group blocks the attack from the si face as shown in Equation 38.80.

–2e–, –H+

S R2NO2S

O

Et4NF-4HF/solvent

S O

R2NO2S

O

F

(38.79)

O R = Me, in MeCN, 40% (100% de) R = Me, in DME, 88% (100% de) R = i-Pr, in DME, 76% (100% de)

S+

S R2N

O

R2N S O O

O

S

F

(38.80)

O O

O

O

F–

They also successfully carried out highly diastereoselective anodic fluorination of N-substituted 1,3-thiazolines and 1,3-oxazolines derived from l-cysteine and l-threonine, respectively [323,324]. Interestingly, the diastereoselectivity increased with an increase of HF content in the supporting fluoride salts and almost 100% de was obtained using Et4NF-5HF as shown in Equation 38.81.

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1497

Stereochemistry of Organic Electrode Processes F

COOMe

COOMe

–2e–, –H+ S

NCOPh

S

F–/DME

(38.81)

NCOPh

Et3N-3HF: 78% (78% de) Et3N-4HF: 78% (94% de) Et4N-5HF: 55% (99% de)

Furthermore, Fuchigami and coworkers investigated anodic fluoro-selenylation and iodofluorination of electron-deficient olefins having chiral auxiliaries [325,326]. However, the diastereoselectivity was low. 7. New Asymmetry Induction Methods Asymmetric electroreduction methods based on new principles have been developed. In recent years, asymmetric synthesis using enzymes has been intensively studied. Enzymatic elecrochemical asymmetric synthesis has also been investigated. Electroenzymatic synthesis is reviewed in Chapter 39. Attachment of enzyme to an electrode surface can be used to catalyze organic redox processes. Sobolov, Fry, and coworkers demonstrated that cross-linked lactate dehydrogenase crystals have high stability and retain their enzymatic activity as biocatalysts during electroenzymatic synthesis of lactic acid from pyruvic acid [327]. Similarly, Kuwabata and coworkers achieved asymmetric synthesis of amino acids like alanine and phenylalanine by cathodic reduction of keto acids using the electrodes on which amino acid oxidase and electron mediator are immobilized [328]. The conversion of glucose to gluconic acid was achieved by inclusion of glucose oxidase in a polypyrrole film [329]. Recently, Schmid and coworkers developed an electrochemical cell for the efficient regeneration of NAD(P)H, which can be coupled to a reduction reaction catalyzed by the thermophilic alcohol dehydrogenase from Thermus sp. as shown in Figure 38.3 [330]. Octane as second organic phase avoided product inhibition and allowed for the production of (1S,3S)-3-methylcycohexanol at a diastereomeric excess of 96% from the corresponding racemic ketone with a productivity of 0.13 g L−1 h−1 and current efficiency of 85%. Quite recently, a highly efficient large-scale asymmetric epoxidation of unfunctionalized olefins was demonstrated using a chiral Mn(III) salen derivative immobilized in a layered crystalline aryldiamine-modified zinc poly(styrene-phenylvinylphosphonate)-phosphate film [331]. Enantiomeric efficiencies up to 99% with 99% conversion were achieved with α-methylstyrene. O

2+

O N

NH2

N Rh

N R NADH

OH2

CH3 rac

Rh-medox TADH +

O N

N Rh

O

NH2

+

+

H Rh-medred

FIgURE 38.3

OH

N R

CH3 NAD+

Reaction scheme of electroenzymatic reaction.

© 2016 by Taylor & Francis Group, LLC

(1S,3S)

CH3 (R)

1498

C.

Organic Electrochemistry

ELECTROCHEMICAL KINETIC RESOLUTION

Electrochemical optical resolution of racemic mixtures is achieved by using enantiomer-differentiating oxidation or reduction. Yamagishi and Aramata [332] also found electrooxidative optical resolution of a racemic Co(1,10-phenanthroline)32+ complex by the chiral clay-coated anode, and Yoshinaga and coworkers [333] optically electroreductively resolved racemic Co(acetylacetonato)3 by using optically active supporting electrolytes. Osa and coworkers [334] reported highly enantioselective electrocatalytic oxidation of racemic monoalcohols using a TEMPO-modified graphite felt electrode in the presence of (−)-sparteine. Almost optically pure R-isomeric alcohols remained unreacted. Highly enantioselective lactonization of racemic diols was also achieved by using the same TEMPO-modified electrode to give (S)-isomeric lactones [335]. On the other hand, Hisaeda and coworkers [336] found that the vitamin B12–mediated electroreductive optical resolution was accompanied by rearrangement: electroreduction of racemic 3-bromo-2-methoxy-2-phenylpropionate using a hydrophobic vitamin B12 mediator afforded ethyl (S)-2-methoxy-2-phenylpropionate in 55% ee, while a novel strapped hydrophobic vitamin B12 provided the corresponding R-enantiomer in 26% ee. Enantioselective electrodialysis of N-α-acetyltryptophans through molecularly imprinted polymeric membranes was also reported [337]. Highly efficient electroenzymatic optical resolution of diols and ketones has been achieved using a polymer-coated graphite felt electrode immobilizing all mediation components including alcohol dehydrogenase [338–340]. Recently, chiral mediators like N-oxyls as shown in Figure 38.4 were demonstrated to be effective for electrochemical asymmetric synthesis and kinetic resolution to provide optically active lactones or alcohols [341–344]. Onomura and coworkers also developed electrochemical asymmetric synthesis using a chiral catalyst as follows. A chiral Lewis acid coordinated with diol or its analogue to activate the O–H bond by chelate formation. Since the complex was readily deprotonated to generate the corresponding alkoxide ion, the complex was more oxidizable than the original alcohol. Successive oxidation provided the corresponding optically active carbonyl compounds as shown in Equation 38.82. R R

OH

Cun+

* R

*

H O Cu

R

YR

– O

* –H+ R

n+

*

Y R

Cun+

*

Y R

R

–2e, –H+

*

Cun+

YR = OH, NHPG, etc.

*

R

O

R

YR

* *

*

O =

O N

N

Ph

Ph

(R,R)-Ph-BOX

NHAc O

Me

Me

Me Me Me

N O

N O

O N

N ArOCO

Me

Cl

Me

BnHNOC

(Ar = 1-naphthyl) [341]

[342]

FIgURE 38.4 Chiral N-oxyl mediators.

© 2016 by Taylor & Francis Group, LLC

[343]

[344]

(38.82)

1499

Stereochemistry of Organic Electrode Processes

Based on this concept, they achieved electrochemical oxidative kinetic resolution of cyclic diols, aminoalcohols, or aminoaldehydes efficiently to obtain optically active compounds, as in Equations 38.83 and 38.84 [345]. OH OH Ph

OH +

OH Ph

OH Ph

Et4NBr/MeOH Cu(OTf )2 (0.1 equiv.) (R,R)-Ph-BOX (0.1 equiv.)

Et4NBr/MeOH Cu(OTf )2 (0.1 equiv.) (R,R)-Ph-BOX (0.1 equiv.)

Bz rac

N

(38.83)

OH Ph 48% (80% ee)

+

CO2Me

OH

N

Bz

(38.84)

Bz (R)

(S)

27% (70% ee)

VI.

OH +

49% (80% ee)

–3e

OH

N

O

–2e

50% (15% ee)

ELECTROCHEMICAL STEREOISOMERIZATION

As shown in Equation 38.85, a stable stereoisomer A can be converted into another stable isomer B through electrochemical oxidation–reduction or reduction–oxidation, if an unstable stereoisomeric intermediate (A±n) is rapidly stereoisomerized to another isomeric intermediate (B±n). Such an electrochemical stereoisomerization sometimes can be easily observed under cyclic voltammetric conditions; for preparative isomerization reactions, AC electrolysis at a suitable frequency or undivided cell electrolysis may be useful. ±ne–

A±n

A

B±n

±ne–

(38.85)

B

Stereoisomerization

A. CIS-TO-TRANS ISOMERIzATION Yeh and Bard [346] observed the isomerization of dialkyl maleates to the fumarates through reduction–oxidation by cyclic voltammetry and rotating ring-disk electrode techniques, as in Equation 38.86. CHCOOR CHCOOR

e–

–

–

CHCOOR CHCOOR

CCHOOR ROOC–CH

e–

CHCOOR ROOC–CH

(38.86)

Stereoisomerization

Electrochemically induced cis-to-trans positional stereoisomerization of ligands has been found on oxidation–reduction of several types of octahedral complexes of group VIIB metals [347–350].

B.

TUB-TO-CHAIR ISOMERIzATION

1,3,5,7-Cyclooctatetraene (COT) ligands in COT–cobalt complexes undergo the tub to chair conformational isomerization induced by electrochemical reduction–oxidation via unstable planar COT− or COT2− [351–355].

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1500

Organic Electrochemistry

C. A-TO-B FORM ISOMERIzATION IN BIANTHRONES The A form bianthrones are isomerized to the B form by both electrochemical reduction–oxidation and oxidation–reduction via unstable anionic and cationic species with bifolded structures, respectively [356–363].

VII.

STEREOISOMERIC EFFECTS ON THERMODyNAMICS AND KINETICS OF ELECTRODE PROCESSES

It is sometimes observed that stereoisomers have different oxidation and reduction potentials and electron-transfer rates. Such phenomena are effectively utilized for clarification of reaction mechanisms and analysis of stereoisomers.

A.

CATHODIC PROCESSES

1. Hydrocarbons Polarographic studies of stereoisomeric α,β-deactivated olefins have been made since early time [364]. Generally, half-wave potentials E1/2 of the first wave of trans-isomers may be slightly less negative than those of the corresponding cis-isomers, although the behaviors of the second waves and limiting current are not simple. Kita and coworkers [365] reported in a kinetic study that the electron-transfer rate of a cis-isomer is several times larger than that of the trans-isomer. Z-isomers of monoactivated olefins may be more easily reduced than the E-isomers in aqueous solutions, but no significant differences are observed in nonaqueous media [366,367]. 2. Halogen Compounds The less stable ax- and exo-monohalo derivatives of monocyclo- and bicycloalkanes, respectively, have remarkably less negative E1/2 values than their eq- and endo-halo isomers. The reduction potentials of vic-dihalides are sensitive to the relative stereochemistry of the halide substituents [368,369]. meso-4,5-Dibromooctane (E1/2 = −1.57 V vs. SCE) is reduced little more easily than its dl-isomer (E1/2 = −1.69 V vs. SCE) [369]. There is also a pronounced dependence of E1/2 on the torsion dihedral angle between the two C-halogen bonds of cyclic vic-dihalides: a plot of E1/2 vs. the dihedral angle ϕ between the two halogen atoms shows minima in E1/2 at ϕ = 0° and 180°, and a maximum at ϕ = 90°. As shown in Figure 38.5, compound 10 having two coplanar bromines is reduced at much less negative potentials compared to its stereoisomer 11 [380]. Such a significant difference in E1/2 values between stereoisomers of vic-dihalides is applicable to stereoisomeric analysis [368–371]. A significant difference in E1/2 was found between cis- and trans-1,2-dihalo olefins [117]. The trans-isomers are more easily reduced. 3. Nitrogen Compounds The E1/2 of cis-azocyclohexane is less negative by about 200 mV at 2 < pH < 9 than that of the trans-isomer, but they are almost equal in aprotic media [372]. On the other hand, aromatic diazo Br t-Bu

Br Br

t-Bu Br 10

E1/2 (V) vs. SCE –1.30

FIgURE 38.5

11 –2.11

Difference in half-wave reduction potentials (E1/2) between stereoisomers of vic-dibromides.

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Stereochemistry of Organic Electrode Processes

1501

compounds show a quite different behavior. Isomeric (cis and trans) azobenzenes are reduced at the almost same potential at pH < 11 because of a fast cis–trans isomerization during the first stage of the reduction, but at pH > 11, the cis-isomer is reduced at a less negative potential than the transisomer [373]. 4. Organometallic Compounds The trans-isomers of (η6-stilbene)Cr(CO)3 and (η6,η6-stilbene)[Cr(CO)3]2 have less negative E1/2 values than the cis-isomers [374]. A similar difference of E1/2 is also observed in the reduction with conformational change of cis- and trans-di-iodocyclohexane [375].

B. ANODIC PROCESSES 1. Hydrocarbons Arnold and Wong [376] have reported that the oxidation potentials of cis- and trans-1,2-diphenylcyclopropanes have simple linear correlations with both their ionization energies and charge transfer absorption energies with tetracyanoethylene as an acceptor. Similar correlations have also been observed in cis- and trans-2,3-diphenyloxiranes. 2. Alcohols Nonaka and coworkers [377] compared the anodic oxidation rates of ax- and eq-hydroxyl groups of a variety of cyclic alcohols. ax-Alcohols with no substituent at the 2-positions are more easily oxidized than the eq-alcohols; for 2-substituted alcohols, the situation is reversed. This is rationalized as due to a different contribution of adsorption between the isomers on an anode. cis-Cyclic diols are much more easily oxidized with chemical oxidizing reagents, such as chromic and periodic acids, than the trans-diols. However, such a significant difference is not observed in the anodic oxidative cleavage [378]. Masui and coworkers [379] reported that threo- and erythroamino alcohols were oxidized at equal potentials to give a slightly different product selectivity. 3. Other Heteroatom Compounds Nelsen and coworkers [380] detected conformational equilibria in eq, eq- and ax, eq-N,N′disubstituted cyclic hydrazines from their oxidation potentials. The anodic oxidation reactions of trans- and cis-1,3-di-isopropyl–2,4-bis(di-isopropylamino)-cyclodiphosph(III)azanes are quite different [381]. The trans-isomer is reversibly oxidized at 0.53 V (vs. SCE) forming a stable cation radical; the cis-isomer undergoes a completely irreversible oxidation at a more positive potential because an unstable radical cation is formed. Evans and coworkers studied structural changes associated with electron-transfer reactions of W(μ5-C5(CH3)5)(CH3)4 and related compounds [382]. The oxidation potential of 2-endo-methylthio-6-endo-substituted norbornanes is strongly influenced by the substituents because of 2,6-space interaction, but there is no such influence in the 6-exo-substituted isomer [383,384]. Yoshida and coworkers found a linear correlation on plotting the oxidation potentials of α-silylated ethers, where the rotation around the C–O bond is restricted, against the HOMO energytorsion angle (Si–C–O–C) curve estimated by MO calculation [385,386].

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Evans, D. H.; Busch, R. W. J. Am. Chem. Soc. 1982, 104, 5057–5062. Evans, D. H.; Xie, N. J. Am. Chem. Soc. 1983, 105, 315–320. Matsue, T.; Evans, D. H. J. Electroanal. Chem. 1984, 163, 137–143. Evans, D. H.; Fitch, A. J. Am. Chem. Soc. 1984, 106, 3039–3041. Wawzonek, S.; Reck, R. C.; Vaught. W. W. Jr.; Fan, J. W. J. Am. Chem. Soc. 1945, 67, 1300–1301. Kita, H.; Nakamura, T.; Itoh, H.; Kano, H. Electrochim. Acta 1978, 23, 405–411. Guillanton, G. L.; Cariou, M. Electrochim. Acta 1977, 22, 619–626. Cariou, M.; Mabon, G.; Guillanton, G. L. Tetrahedron 1983, 39, 1551–1558. Zavada, J.; Krupicka, J.; Sicher, J. Coll. Czech. Chem. Commun. 1963, 28, 1664–1674. Lexa, D.; Saveant, J.-M.; Schäfer, H. J.; Su, K. B.; Vering, B.; Wang, D. L. J. Am. Chem. Soc. 1990, 112, 6162–6177. Klein, J.; Evans, D. H. J. Am. Chem. Soc. 1979, 101, 757–758. O’Connell, K. M.; Evans, D. H. J. Am. Chem. Soc. 1983, 105, 1473–1481. Degrand, C.; Belot, G. Electrochim. Acta 1978, 23, 71–74. Laviron, E.; Mugnier, Y. J. Electroanal. Chem. 1980, 111, 337–344. Rieke, R. D.; Milligan, S. N.; Schulte, L. D. Organometallics 1987, 6, 699–705. Bowyer, J.; Evans, D. H. J. Electroanal. Chem. 1988, 240, 227–237. Arnold, R.; Wong, P. C. Can. J. Chem. 1979, 57, 2098–2012. Nonaka, T.; Abe, K.; Sekine, T.; Denki Kagaku 1979, 47, 184–191; Chem. Abstr. 1979, 47, 184–191. Fuchigami, T.; Nonaka, T.; Watanabe, C.; Yoshiyama, A.; Sekine, T. Denki Kagaku 1983, 51, 812–818; Chem. Abstr. 1984, 100, 164026x. Masui, M.; Kamada, Y.; Sasaki, E.; Ozaki, S. Chem. Pharm. Bull. 1982, 30, 1234–1243. Nelsen, S. F.; Echegoyen, L.; Evans, D. H. J. Am. Chem. Soc. 1975, 97, 3530–3532. Diaz, A. F.; Scherer, O. J.; Andres, K. J. Chem. Soc. Chem. Commun. 1980, 982–983. Lerke, S. A.; Evans, D. H. J. Am. Chem. Soc. 1995, 117, 11768–11772. Wilson, S.; Swanson, D. D.; Klug, J. T.; Glass, R. S.; Ryan, M. D.; Musker, W. K. J. Am. Chem. Soc. 1979, 101, 1040–1042. Glass, R. S.; Coleman, B. R.; Prabhu, U. D. G.; Setzer, W. N.; Wilson, G. S. J. Org. Chem. 1982, 47, 2761–2764. Yoshida, J.; Maekawa, A.; Murata, T.; Matsunaga, S.; Isoe, S. J. Am. Chem. Soc. 1990, 112, 1962–1970. Yoshida, J. In: Topics in Current Chemistry, 170, Electrochemistry. (Steckhan, E. ed.) Springer-Verlag, Berlin, Germany, 1994, pp. 39–82.

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39

Electroenzymatic Synthesis* Christina Kohlmann and Stephan Lütz

CONTENTS I. Introduction ........................................................................................................................1511 II. NAD(P)+-Dependent Electroenzymatic Oxidations...........................................................1514 III. NAD(P)H-Dependent Electroenzymatic Reductions .........................................................1519 IV. Flavin-Dependent Oxidations with Oxidative Regeneration ............................................. 1528 V. Flavin-Dependent Reactions with Reductive Regeneration ............................................... 1531 VI. Electroenzymatic Oxidations with Peroxidases .................................................................1533 VII. Synthesis Reactions with Electrochemical Substrates Supply or Product Conversion ......1536 VIII. Summary and Outlook ....................................................................................................... 1538 References .................................................................................................................................... 1539

I. INTRODUCTION The use of biocatalysts, that is, whole cells or isolated enzymes in organic synthesis, has been constantly growing over the last decades, both on lab scale [1–3] and industrial scale [4,5]. Enzymes show remarkable chemo-, regio-, and stereoselectivities and are therefore used in a broad range of resolution reactions or asymmetric syntheses, for example, in the synthesis of chiral building blocks (amines, alcohols), carbohydrate and peptide chemistry, or other transformations. Especially the improvements in molecular biology during the last decade, particularly in the field of protein engineering, allowed to create biocatalysts with properties adapted to chemical synthesis, for example, higher selectivity or process stability [6–8]. This has lead to an overall increase in the application of enzymes for chemical purposes. Enzymes are classified according to the reactions they catalyze. For a combination of enzymes and electrochemistry, mainly enzymes catalyzing redox reactions are of interest, these form the class of oxidoreductases, the first class according to the enzyme nomenclature (E.C. 1.X.X.X) [9]. In contrast to reactions catalyzed by hydrolases, which are mostly resolution of racemic compounds, biocatalytic redox reactions in organic chemistry are typically asymmetric syntheses, converting prochiral substrates into chiral products. Some typical oxidoreductases and some of their reactions are summarized in the following list: 1. Dehydrogenases a. Alcohol dehydrogenases (ADH)/ketoreductases (KRED): Reduction of prochiral ketones to chiral alcohols [10,11], resolution of racemic alcohols by selective oxidation, formation of lactones from meso-diols b. Amino acid dehydrogenases (AADH): Reduction of α-keto acids to chiral amino acids c. Polyol dehydrogenases (glycerol DH, sorbitol DH, mannitol DH, aldose reductases): Carbohydrate reduction, polyol resolution d. Ene reductases/enoate reductases: Enantioselective hydrogenation of unsaturated enoates or similar compounds * In memoriam Prof. Dr. E. Steckhan (1943–2000).

1511

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1512

Organic Electrochemistry

2. Monooxygenases: regio- and stereoselective hydroxylations, epoxidations, S-,N-oxidations, Baeyer–Villiger oxidations 3. Oxidases: egio- and stereoselective oxidation of carbohydrates, hydroxylation of phenols, oxidation of amino acids to α-keto acids 4. Peroxidases: Epoxidations, S-,N-oxidations While using oxidoreductases for synthesis reactions, there is always a need for a source or sink for the electrons involved, that is, a redox equivalent. Nature has developed a quite diverse molecular repertoire to achieve this transfer of redox equivalents, for example, between protein domains or between proteins and metabolites (Figure 39.1). This function is fulfilled by freely dissociated cofactors as well as enzyme-bound coenzymes or prosthetic groups. From this selection of cofactors, three are mainly important in redox biocatalysis on preparative scale: (1) the nicotinamides (NAD(P)/H), for example, for ADHs or monooxygenases; (2) the flavins (FAD/H2), for example, for oxidases or monooxygenases, and (3) heme, for example, in P450-monooxygenases (CYPs) or heme peroxidases. As these cofactors/coenzymes are relatively expensive [12], they cannot be used as stoichiometric redox equivalents. Thus there is a need for methods to regenerate the desired redox state of the cofactors involved in order to make reactions with oxidoreductases economically feasible. In addition, the cofactor in the contrary redox state might be inhibitory, so constant adjustment of the right redox state is essential for a successful application of oxidoreductases in organic synthesis. Several efficient methods have been developed for cofactor recycling [13], which include enzymatic and various chemical methods (transition-metal catalyzed [14], photochemical, electrochemical [15]) as well as combined methods. The principle of the different systems is illustrated for a nicotine-amide-dependent enzyme (Figure 39.2). In enzyme-coupled regeneration methods, a second enzyme is applied that converts a sacrificial cosubstrate to a coproduct while converting the cofactor back into the desired form [16]. Another approach is called substrate-coupled cofactor regeneration, in cases where the same enzyme is used both for the production and  regeneration reaction, for example, dehydrogenases that can oxidize isopropanol to acetone [11,17,18]. For flavin-dependent oxidases, regeneration can,  for example, be done with molecular oxygen, which leads to the formation of hydrogen peroxide as coproduct, a substance that is known to inactivate various enzymes. The hydrogen peroxide formed can be destroyed by applying catalase, but then again a second enzyme is required. The most typical regeneration approaches are summarized in Table 39.1. It is worth noting that so far only the enzymatic regeneration systems have been applied on industrial scale. Electrochemistry in particular is a promising tool for the regeneration of the different cofactors and coenzymes. For this technology, an oxidoreductase-catalyzed biotransformation coupled to an electrochemical supply of redox equivalents, the term electroenzymatic synthesis has been established [19–23]. The appeal of this combination lies in the fact that the electron allows a quasi mass-free regeneration of the cofactor and thus avoids the coupled coproduct. Some drawbacks have to be considered nevertheless: specially designed bioreactors (i.e., electrochemical cells) are required for the synthesis reactions. The combination of isolated enzymes with electrochemical reactions is well established in biosensor techniques for analytical purposes. In electrochemical biosensors, the biocatalyst is usually bound onto the electrode surface. While this arrangement is useful for analytical applications, it is in most cases not attractive for synthetic applications. It is lacking both in longterm stability and  productivity due to the low current densities. For preparative scale applications, as in electroenzymatic synthesis, the redox transfer between electrode surface and protein (typically homogeneously solubilized in the aqueous electrolyte) is therefore often carried out via a small molecule mediator in the form of an indirect electrolysis [24]. This chapter gives an overview of enzymatic synthesis reactions where the cofactors or coenzymes involved are regenerated electrochemically or the required cosubstrates of an enzymatic reaction are generated electrochemically. Its focus is on preparative application rather than mechanistic or

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1513

Electroenzymatic Synthesis NH2 N H2C

O

O

O P

O P

HO CH

N

O CH2 O N H H OH H H OH OH

OH

HO CH

N

HOOC NH2

HO CH CH2 N

S

NH2

O

N

HOOC OH

NH N O flavin adenine dinucleotide (FAD) NH2

thio-tyrosine O

N

N

– – O O O P O P O O O

N

N

O RO

OH

O HO

NH2

+ N

OH

NAD+ (R=H) NADP+ (R=PO32–)

O

H N

H N

C

O C

O O

CH2

NH H N

COOH

N

O

HO NH O

O

COOH C

HOOC O

O

6-hydroxy-DOPA methoxanthine (PQQ) O

N H tryptophan tryptophylquinone (TTQ)

OH N

N Fe N

N OH O

heme

FIgURE 39.1

Cofactors and prosthetic groups.

analytical studies. Since this type of interdisciplinary research is published in a variety of journals covering enzymology, electrochemistry, and organic synthesis, no common reporting system of the data can be realized (e.g., biotechnology journals do not necessarily report current yields, while electrochemistry journals typically do not report cycle numbers for the protein). Wherever possible, standardized units are given to make a comparison of procedures easier. Important values to compare enzymatic processes are the enzyme and cofactor utilization, usually expressed as total

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1514

Organic Electrochemistry Enzyme coupled E1

O R´

Substrate coupled O

OH

R˝

R´

R˝

NAD(P)H NAD(P)+ Coproduct

E2

Cosubstrate

R´

R´

R˝ NAD(P)H

Coproduct

E1

Chemical OH

E1

R˝

O R´

NAD(P)+ Cosubstrate

OH

E1 R´

R˝ NAD(P)H

Coproduct

R˝

NAD(P)+

Cat

Cosubstrate

FIgURE 39.2 Principle of cofactor regeneration shown for an NAD(P)-dependent dehydrogenase reaction as example (E1, production enzyme; E2, regeneration enzyme; Cat, catalyst).

TAbLE 39.1 Selected Cofactor Regeneration Systems Type Enzymatic

Chemical

Catalyst

Cofactor

Formate dehydrogenase (FDH) Glucose dehydrogenase (GDH) Glucose-6-phosphate dehydrogenase (G6PDH) Phosphite Dehydrogenase (PDH) Hydrogenase (H2ase) Alcohol Dehydrogenase (ADH) O2/Catalase O2/Glucose Oxidase Rhodium bipyridine (bpy) complexes Phenanthroline dione (PD) derivatives 2,2′-azinobis(3-ethylbenzothiazoline)-6-sulfonate (ABTS) Ferrocene carboxylic acid and derivatives

NAD(P)H NADPH NADPH NADPH NADPH NAD(P)H FAD+ Heme/Compound I NAD(P)H, FADH2 NAD+, NADP+ NAD+ FAD+

turnover number (ttn), referring to the amount of product obtained per amount of (co)enzyme used up in the reaction (ttn = n(product)/n((co)enzyme)). Another important value is the unit (U), which is a measure for catalytic activity. One unit is defined as the enzymatic activity to convert 1 µmol of substrate per minute (i.e., a chemical reaction rate). For further details on enzymatic processes and their parameters, the reader is referred to the literature, for example [4].

II.

NAD(P)+-DEPENDENT ELECTROENZyMATIC OXIDATIONS

Syntheses where the cofactor NAD(P)+ has to be regenerated to its oxidized form can be carried out with direct, indirect, and enzyme-coupled electrochemical cofactor regeneration. In the latter case, a second enzyme is used as the electron shuttle. For the regeneration of NAD(P)+ directly at an anode, high overpotentials of at least 900 mV versus the saturated calomel electrode (SCE) [19] are required. Therefore, this method can only be used for reactions with compounds that do not undergo reactions themselves under these conditions. The direct regeneration of NAD+ is employed in the synthesis of d-gluconolactone from d-glucose by glucose dehydrogenase (Table 39.2, entry 1) [25] in a plug flow reactor. In this setup, the working and auxiliary compartments were separated by a membrane. In this experiment, oxidation of NADH at an imposed potential lower than 0 mV versus SCE was reported. A product formation of 13 g L−1 was possible at a potential of 700 mV versus SCE. Steckhan and coworkers used different complexes of phenanthrolinedione as mediators (see Figure 39.3). 4,10-Dioxa-tricyclo[5.2.1.02,6]dec-8-en-3-one was synthesized from meso-5,6-dihydroxy-methyl7-oxabicyclo[2.2.1]hept-2-ene by an ADH using 1 µmol of N-methyl-1,10-phenanthrolinium-5,6-dione

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Entry 1 2

3 4 5

6

7

8 9

Enzyme glucose dehydrogenase (GDH, E.C. 1.1.99.10) ADH from horse liver (HLADH, E.C. 1.1.1.1) ADH from horse liver (HLADH, E.C. 1.1.1.1) ADH from horse liver (HLADH, E.C. 1.1.1.1) ADH from yeast (YADH, E.C. 1.1.1.1) ADH from Thermoanaerobium brockii (TBADH, E.C. 1.1.1.1) glycerol dehydrogenase from Cellulomonas sp. (GDH, E.C. 1.1.1.6) ADH from horse liver (HLADH, E.C. 1.1.1.1.) glucose dehydrogenase (GDH, E.C. 1.1.1.47)

Substrate

Product

Cofactor

Mediator

Experimental Data

Literature

d-glucose

d-glucono-δ-lactone

NAD+/H

None

13 g L−1, ttncofactor > 10,000 Conversion 74%, tofmediator 30 h−1

[25]

meso-5,6-dihydroxymethyl-7-oxabicyclo[2.2.1]hept-2-ene cyclohexanol

4,10-dioxa-tricyclo[5.2.1.02,6]dec-8-en-3-one

NAD+/H

PDMe (BF4)2

cyclohexanone

NAD+/H

meso-3,4-dihydroxymethylcyclohexene 2-hexen-1-ol

3a,4,7,7a-tetrahydro-3Hisobenzofuran-1-one 2-hexenal

NAD+/H

Ru(PD)3 (ClO4−)2 PDMe (BF4)2

Conversion 75%

[27]

Conversion 99.5%, tofmediator 35 h−1 1.77 mM, ttncofactor 18, ttnmediator 36, current efficiency 90 4.1 mM, ttncofactor 41, ttnmediator 82, current efficiency 95% 30 mM, reaction time 240 h

[20,26]

NAD+/H

Fe(tmphen)3

2-butanol

2-butanone

NADP+/H

Fe(tmphen)3

rac-phenylethane-1,2diol

(S)-phenylethane-1,2-diol

NAD+/H

ABTS

meso-3,4-dihydroxymethylcyclohexene d-glucose

3a,4,7,7a-tetrahydro-3Hisobenzofurane-1-one gluconic acid

NAD+/H

ABTS

Yield 93.5%, ee > 99.5

[30]

NAD+/H

3,4-dihydroxybenzaldehyde

ttncofactor 100

[31]

Electroenzymatic Synthesis

TAbLE 39.2 NAD(P)+-Dependent Electroenzymatic Oxidations

[26]

[28]

[28]

[29]

(Continued)

1515

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1516

TAbLE 39.2 (Continued) NAD(P)+-Dependent Electroenzymatic Oxidations Entry 10

11 12 13 14 15

Enzyme

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Product

Cofactor

6-phosphogluconate

Substrate

ribulose 5-phosphate

NADP+/H

CAV, AMAPOR

Mediator

Conversion 98% in 2.5 h

[32]

(2R,3S)-isocitrate

2-oxoglutarate

NADP+/H

2-oxoglutarate

NADP+/H

cyclohexanol

cyclohexanone

NAD+/H

cyclohexanol

cyclohexanone

NAD+/H

13,000 mmol * kg−1 * h−1 after 3.8 h (ee > 99%) 14,000 mmol*kg−1*h−1 after 3.6 h (ee > 99%) 2 mM, quantitative oxidation within 10 h Not available

[32]

(2R,3S)-isocitrate

CAV, AMAPOR A-Q-2-S, AMAPOR Pd derivative

d,l-glyceraldehyde

d-glyceric acid

NAD+/H

ferrocene + diaphorase ABTS

Experimental Data

8.6 g L−1 d−1 ee 88%

Literature

[32] [33] [34] [35]

Organic Electrochemistry

6-phosphogluconate dehydrogenase from pig heart (E.C. 1.1.1.44) isocitrate dehydrogenase from pig heart (E.C. 1.1.1.41) isocitrate dehydrogenase from pig heart (E.C. 1.1.1.41) ADH from horse liver (HLADH, E.C. 1.1.1.1.) ADH from horse liver (HLADH, E.C. 1.1.1.1.) aldehyde dehydrogenase from Deinococcus geothermalis

1517

Electroenzymatic Synthesis O O O

O

O

N+

N

O N O

N 2+

N

M

N

N

N X2

N

Ru2+

N

O

(ClO4)2

N N

CH3

N N O O

Heteroleptic M(TPA)(PD) complexes (M = Ru, Co)

FIgURE 39.3

Homoleptic [Ru(PD)3](ClO4)2 complex

PDMe+(BF4)–

Phenanthroline dione (PD) complexes and derivatives for NAD(P)+ regeneration.

tetrafluoroborate (PDMe(BF4)2) as mediator leading to 74% conversion and 30 turnovers per hour for the mediator (Table 39.2, entry 2) [26]. A ruthenium complex of phenanthrolinedione (Ru(PD)3(ClO4)2) was applied in the conversion of cyclohexanol to cyclohexanone. With 25 U of ADH from horse liver (HLADH), 0.5 mM NAD+, 0.1 mM mediator, 10 mM substrate and a potential of 100 mV versus Ag|AgCl 75% conversion was achieved (Table 39.2, entry 3) [27]. The mediators PDMe2+ and Ru(PD)32+ were also used in the synthesis of tetrahydro-3H-isobenzofuranone (Table 39.1, entry 4) [20,26] leading to a turnover number of 35 h−1 for the mediator (see Figure 39.4). Iron 3,4,7,8-tetramethylphenanthroline (Fe(tmphen)3) is also suitable for the regeneration of NAD+. It was used as mediator in the synthesis of 2-hexenal catalyzed by ADH from yeast (Table 39.2; entry 5) [28] and 2-butenal by T. brockii ADH (Table 39.2, entry 6) [28]. The electrolysis was conducted in a divided cell with a potential of 0.63 V versus Ag|AgCl, and current efficiencies of ~90% were obtained. For the synthesis of 2-hexenal, the total ttn of the cofactor was 18 and for the mediator 36. In the synthesis of 2-butenal, ttns of 41 for NAD+ and 82 for Fe(tmphen)3 were achieved. The mediator 2,2′-azinobis(3-ethylbenzothiazoline)-6-sulfonate (ABTS) was used in the synthesis of (S)-phenylethane-1,2-diol (see Figure 39.5) by glycerol dehydrogenase (GDH) (Table 39.2, entry 7) [29] as well as tetrahydro-3H-isobenzofuranone by HLADH (Table 39.2, entry 8) [30]. The oxidative resolution of racemic phenylethane-1,2-diol was carried out in an electrochemical

NAD+

PDMediatorred

NADH

PDMediatorox

Anode

OH OH

HLADH

O

2e–

O

FIgURE 39.4 Synthesis of lactones by electroenzymatic oxidation with horse liver alcohol dehydrogenase (HLADH) and PDMe+ as mediator.

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1518

Organic Electrochemistry OH

O OH

ABTS2–

NAD+

acyloine

S-dial

Anode

OH +

GDH OH ABTS–

NADH

OH rac-dial

FIgURE 39.5 mediator.

Oxidative resolution of racemic diols with glycerol dehydrogenase (GDH) and ABTS as

membrane reactor yielding enantiopure (S)-diol. The oxidation products are extracted into an organic phase to overcome the inhibition of the enzyme. In contrast, the synthesis of tetrahydro-3Hisobenzofuranone was performed in a quasi-divided electrochemical cell leading to 93.5% conversion yielding the product in an optical purity of ee > 99.5%. A reactor with glucose dehydrogenase and 3,4-dihydroxybenzaldehyde as mediator immobilized on a carbon felt electrode for the synthesis of gluconic acid was reported by Manjon et al. (Table 39.2, entry 9) [31]. A divided cell reactor was used and operated at 0.7 V versus Ag|AgCl and also at 0.2 V versus Ag|AgCl. As expected, higher potentials led to faster rate, while the ttn was about 100 for the cofactor in both experiments. The oxidative decarboxylation of 6-phosphogluconate to ribulose 5-phosphate with 6-phosphogluconate dehydrogenase was carried out using the mediator carboxamidomethyl viologen (CAV) in combination with an artificial mediator accepting pyridine nucleotide oxidoreductase (AMAPOR) (Table 39.2, entry 10) [32]. This is an example where the indirect regeneration of NAD(P)+ is coupled with an enzymatic reaction step. In a reaction containing 20 mM substrate, 3 mM CAV, 0.5 mM NADP+, 15 U of 6-phosphogluconate-dehydrogenase, and 20 U AMAPOR, a yield of 98% was obtained in 2.5 h. The applied potential was −200 mV versus SCE. The resolution of racemic isocitrate via oxidative decarboxylation leads to enantiopure (2R,3S)-isocitrate and 2-oxoglutarate. This reaction can be catalyzed by an isocitrate dehydrogenase. The required cofactor NADP+ was regenerated using AMAPOR and either CAV (Table 39.2, entry 11) [32] or anthraquinone 2-sulphonate AQ-2-S (Table 39.2, entry 12) [32] as mediator. Both approaches lead to comparable results; for CAV space time yields of 13,000 mmol product per kg catalyst and hour and for AQ-2-S space time yields of 14,000 mmol product per kg catalyst and hour were obtained. For substrate concentrations of up to 400 mM of the racemate, full conversion of (2S,3R)-isocitrate was achieved leading to an ee of higher than 99% for (2R,3S)-isocitrate. The conversion of cyclohexanol into cyclohexanone by HLADH was combined to a regeneration approach using quinone and diaphorase (Table 39.2, entry 13) [33]. Full conversion of the substrate was achieved in a setup containing 2 mmol cyclohexanol, 0.1 mM NAD+, 0.5 µmol HLADH, 0.1 mM mediator, and 10 U mL−1 of diaphorase. The electrolysis was carried out using glassy carbon working electrodes and a potential of −100 mV versus SCE. Oxidation of cyclohexanol to cyclohexanone was also performed by using a ferrocene/ diaphorase/HLADH immobilized electrode (Table 39.2, entry 14) [34]. For preparing this electrode, aminoferrocene and 2-aminoethylferrocene were bound to a graphite felt coated with a thin polyacrylic acid film. Subsequently, the electrode was treated with diaphorase and finally with ADH. After this preparation, the electrode was successfully used in the oxidation of NADH to NAD+. Recently, aldehyde dehydrogenase (ALDH) from D. geothermalis DSM 11300 was used for the oxidative resolution of glyceraldehyde (Table 39.2, entry 15) [35].

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1519

Electroenzymatic Synthesis

III.

NAD(P)H-DEPENDENT ELECTROENZyMATIC REDUCTIONS

The reductive regeneration of NADP(H) from NAD(P)+ is more delicate than the reverse reaction, since it effectively requires a selective hydride addition to yield the enzymatically active form of the cofactor. The direct electrochemical reduction of NAD(P)+ is a two-step reaction; for the first step, a potential of approximately −1.2 V versus SCE is necessary to transfer one electron to NAD(P)+ forming a radical species [36]. A second electron can be transferred unto this radical after protonation (potential varies from −1.7 to −2.0 V vs. SCE) [36–39]. The radicals formed in the first reduction step can undergo side reactions, for example, dimerization, leading to enzymatically inactive forms of the cofactor. Moreover, protonation is not selective; besides the desired 1,4-NAD(P)H, enzymatically inactive 1,6-NAD(P)H is likely to occur. After a few cycles, no enzyme-active 1,4-NAD(P)H is left, so that direct electrochemical reduction is not feasible. Only two examples of direct regeneration of NADH are reported. In these cases, very low concentrations of cofactor are used. It is assumed that with low concentrations, the dimerization is negligible. One example is the reduction of pyruvate to d-lactate by d-lactate dehydrogenase (Table 39.3, entry 1) [40]. Cofactor regeneration takes place at a cholesterol-modified gold amalgam electrode. With this setup, a turnover number for NADH of ~1400 and a conversion of ~72% was achieved. The other example deals with the synthesis of l-glutamate from 2-oxoglutarate by an l-glutamate dehydrogenase (Table 39.3, entry 2) [41]. Vanadium-silica gels were added to the electrolyte to increase the conductivity. Complete conversion of the substrate and a total turnover number ttn of 3300 for the cofactor were obtained. Steckhan has therefore suggested the following criteria for a mediator to enable indirect electrochemical reduction of NAD(P)+ to NAD(P)H [19]: 1. The mediator must transfer two electrons or one hydride ion in one step. 2. The electrochemical activation of the mediator must be possible at potentials less negative than −0.9 V versus SCE (at more negative potentials the direct reduction of NAD(P)+ takes place). 3. The mediator must not transfer the electrons to the substrate. 4. Only enzymatically active 1,4-NAD(P)H must be formed. One of the first substances that met all four requirements for selective NAD(P)+ reduction was the (2,2′-bipyridyl)rhodium complex (Rh(bpy)). One example where this mediator is used to regenerate NADH is the reduction of cyclohexanone to cyclohexanol catalyzed by an ADH (Table  39.3, entry 3) [21]. 0.25 mmol of [Rh(bpy)3]2+, 0.1 mmol of NAD+ and 1.12 mmol cyclohexanol were dissolved in a tris(hydroxymethyl)aminomethane (TRIS) buffer (pH 9.0). With this system, 2.9 cycles were achieved for NAD regeneration and 1.2 cycles for mediator regeneration. The small number of regeneration cycles is assumed to result from a passivation of the cathode by a layer of [Rh(bpy2) (H2O)2]Cl or [Rh(bpy2)(OH)2]Cl formed on the electrode surface. The next generation of rhodium mediators was developed by using the pentamethycyclopentadienyl anion (Cp*) as ligand. With this mediator, 1,4-NAD(P)H is formed with a regioselectivity of more than 99%. The reaction equilibria involved are summarized in the following, adapted from [42]:

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[Cp*Rh III (bpy)L]2+ + e − ⇌ [Cp*Rh II (bpy)L]+

(39.1)

[Cp*Rh II (bpy)L]+ ⇌ [Cp*Rh II (bpy)]+ + L

(39.2)

[Cp*Rh II (bpy)L]+ + e − ⇌ Cp*Rh I (bpy) + L

(39.3)

1520

TAbLE 39.3 NAD(P)H-Dependent Electroenzymatic Reductions Entry 1 2

3 4

5 6 7 8

10

11

d-lactate dehydrogenase (d-LDH, E.C. 1.1.1.28) l-glutamate dehydrogenase from bovine liver (GDH, E.C. 1.4.1.3) ADH from horse liver (HLADH, E.C. 1.1.1.1) d-lactate dehydrogenase from Staphylococcus epidermidis (d-LDH, E.C. 1.1.1.28) ADH from horse liver (HLADH, E.C. 1.1.1.1) ADH from Rhodococcus sp. (S-ADH, E.C. 1.1.1.2) ADH from horse liver (HLADH, E.C. 1.1.1.1) ADH from horse liver (HLADH, E.C. 1.1.1.1) ADH from Lactobacillus brevis (LbADH, E.C. 1.1.1.1) ADH from L. brevis (LbADH, E.C. 1.1.1.1) ADH from L. brevis (LbADH, E.C. 1.1.1.1)

Substrate

Product

Cofactor

Mediator

pyruvate

d-lactate

NAD /H

None

2-oxoglutarate

l-glutamate

NAD+/H

None

cyclohexanone

cyclohexanol

NAD+/H

[Rh(bpy)3]2+

pyruvate

d-lactate

NAD+/H

[Cp*Rh(bpy)Cl]+

4-phenyl-2-butanone

NAD+/H

cyclohexanone

(S)-4-phenyl-2butanol (S)-4-phenyl-2butanol cyclohexanol

cyclohexanone

Experimental Data

Literature

Conversion 72%, 18.2 mM in 21 h, ttncofactor 1400 Conversion 100%, ttncofactor 3300

[40]

Conversion ~ 26%, ttncofactor 2.9, ttnmediator 1.2 14 mM, ee 93.5%, ttncofactor 14, ttnmediator 7, current efficiency 67%

[21]

Cp*Rh(bpy)L

Conversion 70%, ee 65% (S-product)

[19]

NAD+/H

Cp*Rh(bpy)L

Conversion 76%, ee 77% (S-product)

[19]

NAD+/H

Cp*Rh(bpy)L

Conversion 92% in 42 days

[44]

cyclohexanol

NAD+/H

Cp*Rh(bpy)L

Conversion 100% in 3 days

[44,45]

acetophenone

(R)-phenylethanol

NADP+/H

Cp*Rh(bpy)L

[46]

acetophenone

(R)-phenylethanol

NADP+/H

Cp*Rh(bpy)L

acetophenone

(R)-phenylethanol

NADP+/H

Cp*Rh(bpy)L

Conversion 98%, ee > 99.9%, 13.5 g L−1 d−1, ttnenzyme 75,000, ttncofactor 35, ttnmediator 35, current efficiency 55% Conversion 93%, ee > 98%, 9.0 g L−1 d−1, ttnenzyme 21,000, ttncofactor 12, ttnmediator 55, current efficiency 93% Conversion 98%, ee > 99.9%, 2.6 g L−1 d−1, ttnenzyme 5000, ttncofactor 64, ttnmediator 64 current efficiency 45%

4-phenyl-2-butanone

+

[41]

[43]

[46]

[46]

(Continued)

© 2016 by Taylor & Francis Group, LLC

Organic Electrochemistry

9

Enzyme

Entry

Enzyme

Substrate

Product

Cofactor

Mediator

Experimental Data

12

ADH from L. brevis (LbADH, E.C. 1.1.1.1)

4-chloroacetophenone

(R)-4chlorophenylethanol

NADP+/H

Conversion 70%, ee > 97.3%, ttnmediator 200

[51]

13

ADH from Thermus sp.

NAD+/H

MV2+, diaphorase

15

ADH (ADH, E.C. 1.1.1.1)

MV2+, diaphorase

ADH (ADH, E.C. 1.1.1.1)

NAD+/H

MV2+, diaphorase

17

d-lactate dehydrogenase from Leuconostoc mesenteroides (d-LDH, E.C. 1.1.1.28) benzoylformate dehydrogenase from Enterococcus faecalis (BFR) benzoylformate dehydrogenase from Enterococcus faecalis (BFR) d-lactate dehydrogenase (d-LDH, E.C. 1.1.1.28)

(1S,2S)-(+)-2methylcyclohexanol (1S,3S)-(-)-2methylcyclohexanol d-lactate

NAD+/H

16

2-methylcyclohexanone rac-3-methylcyclohexanone pyruvate

NAD+/H

MV2+, diaphorase

3.1 g L−1 d−1 (for the organic phase) de > 96%, current efficiency 85% Conversion 100%, current efficiency 97.8 Conversion 48.9%, ee > 99%, ttnmediator 91, current efficiency 98.6% Conversion 51.7%, ee 93.1%, ttnmediator 94, current efficiency 96.5% Conversion 80%

[52]

ADH (ADH, E.C. 1.1.1.1)

(1S,3S)-3methylhexanol cyclohexanol

NADP+/H

14

rac-3methylcyclohexanone cyclohexanone

Cp*Rh(bpy)L (polymerbound) Cp*Rh(bpy)L

benzoylformate

(R)-mandelate

NAD+/H

MV2+, diaphorase

40 mM in 30 h, ttncofactor 133

[55]

benzoylformate

(R)-mandelate

NAD+/H

FAD + LipDH

47.5 mM in 18 h, ttncofactor 158

[55]

pyruvate

d-lactate

NAD+/H

MV2+ + LipDH

[56]

cyclohexanone

cyclohexanol

NAD+/H

MV2+ + LipDH

Conversion 94% in 9 days, ee 94%, ttnenzyme 3.5·107, ttncofactor 940, ttnLipDH 5.4·105, current efficiency 104 ± 10% Conversion 65%

18

19

20

21

ADH from horse liver (HLADH, E.C. 1.1.1.1)

Literature

[53]

Electroenzymatic Synthesis

TAbLE 39.3 (Continued ) NAD(P)H-Dependent Electroenzymatic Reductions

[53] [53] [54]

[45] (Continued)

1521

© 2016 by Taylor & Francis Group, LLC

1522

TAbLE 39.3 (Continued ) NAD(P)H-Dependent Electroenzymatic Reductions Entry

Enzyme

Substrate

Product

Cofactor

Mediator

Experimental Data

Literature [56]

22

d-lactate dehydrogenase (d-LDH, E.C. 1.1.1.29)

pyruvate

d-lactate

NAD /H

MV + FDR

23

glutamate dehydrogenase (GluDH, E.C. 1.4.1.3)

2-oxoglutarate

l-glutamate

NADP+/H

MV2+ + FDR

24

glutamate dehydrogenase from beef liver (GluDH, E.C. 1.4.1.3) glutamate dehydrogenase from beef liver (GluDH, E.C. 1.4.1.3) glutamate dehydrogenase from beef liver (GluDH, E.C. 1.4.1.3) ADH from T. brockii or equine liver (ADH, E.C. 1.1.1.1 or E.C. 1.1.1.2) ADH from T. brockii or equine liver (ADH, E.C. 1.1.1.1 or E.C. 1.1.1.2) l-glutamate dehydrogenase from bovine liver (GDH, E.C. 1.4.1.2)

2-oxoglutarate

l-glutamate

NADP+/H

MV2+, AMAPOR

Conversion 90% in 14 days, ee 94%, ttnenzyme 2.2·107, ttncofactor 900, ttnFDR 7.3·106, current efficiency 103 ± 10% Conversion 100% in 7 days, ttnenzyme 1.1·107, ttncofactor 1000, ttnFDR 7.5·106, current efficiency 105% ± 10% 29 mol*kg−1*h−1, ttncofactor 29,000

2-oxoglutarate

l-glutamate

NAD+/H

MV2+, AMAPOR

7 mol*kg−1*h−1, ttncofactor 29,000

[32]

2-oxoglutarate

l-glutamate

NADP+/H

CoSep, AMAPOR

1 mol*kg−1*h−1

[32]

Various ketones and alcohols

Alcohols

acetophenone

Current efficiencies 76–91%, 0.3–4.7 µmol of product

[57]

Various ketones and alcohols

Alcohols

[57]

l-glutamate

MV2+ + FNR or MV2+ + diaphorase hydrogenase

Current efficiencies 89–100%, 3.2–18.2 µmol of product

2-oxoglutarate

NADP+/H or NAD+/H NADP+/H or NAD+/H NAD+/H

ttncofactor 992–1034, 99.2–103.4 mmol L−1 in 5 h

[58]

25

26

27

29

© 2016 by Taylor & Francis Group, LLC

2+

[56]

[32]

Organic Electrochemistry

28

+

1523

Electroenzymatic Synthesis

[Cp*Rh II (bpy)]+ + e − ⇌ Cp*Rh I (bpy)

(39.4)

[Cp*Rh II (bpy)L]+ + [Cp*Rh II (bpy)]+ ⇌ Cp*Rh I (bpy) + [Cp*Rh III (bpy)L]2+

(39.5)

Cp*Rh I (bpy) + H + ⇌ [Cp*Rh I (bpy)H]+

(39.6)

[Cp*Rh I (bpy)H]+ + NAD(P)+ + L ⇌ [Cp*Rh III (bpy)L]2+ + NAD(P)H

(39.7)

The first synthetic application was the conversion of pyruvate to d-lactate by d-lactate dehydrogenase (Table 39.3, entry 4) [43]. A solution of 1 mM [Cp(Me)5Rh(bpy)Cl]Cl, 2 mM NAD+, 20 mM substrate and 1300 U enzyme was electrolyzed at −0.6 V vs. Ag|AgCl to yield 14 mM d-lactate (ee 93.5%) and turnover numbers of 14 for the mediator and 7 for the cofactor. This type of mediator has found widespread applications and its mechanism is widely studied [19,20,42]. Steckhan and coworkers investigated the reduction of 4-phenyl-2-butanone to (S)-4phenyl-2-butanol by HLADH (Table 39.3, entry 5) [19] and by an ADH from Rhodococcus sp. (Table 39.3, entry 6) [19] As mediator Cp*Rh(4-ethoxy-methyl-2,2′-bpy)L and several types of water-soluble polymer-bound versions of the mediator were used and applied in batch reactors as well as in an electrochemical enzyme membrane reactor (EEMR) (Figure 39.6). Bergel and coworkers used the mediator [Cp(Me)5Rh(bpy)Cl]Cl in a setup called a dialysis membrane electrochemical reactor (D-MER, with a membrane consisting of regenerated cellulose) (Table 39.3, entry 7) [44] and in an ultrafiltration dialysis membrane electrochemical reactor (UF-MER, with membrane made of cellulose) (Table 39.3, entry 8) [44,45], to confine the catalyst to the carbon working electrode compartment. By using these two different reactors and 100 mM cyclohexanone as substrate, a conversion of 92% in 42 days, or 100% in 3 days, was achieved. For the D-MER setup, 0.5 mM rhodium complex, 1 mM NAD+, and 23 U of HLADH were applied and in the UF-MER approach 1 mM Rh complex, 1 mM NAD+, and 73 U of HLADH. Hildebrand and Lütz reported the synthesis of (R)-phenylethanol catalyzed by an ADH from L. brevis with Cp*Rh(bpy) as mediator (Table 39.3, entry 9) [46]. Excellent ee’s of over 99.9% as well as high space–time yields of 14 g L−1 d−1 were reported. Reactions were conducted in a 200 mL batch reactor using carbon felt as working electrode and a platinum net in a dialysis sack as the counter electrode. Enzyme stabilization by the addition of bovine serum albumin yielded total conversion and total turnover numbers of ttn (ADH) = 75,000. Cofactor and mediator showed total turnover numbers of ttn = 35. The application of an immobilized enzyme preparation [18] to the same reactor led to increased turnover numbers for the mediator (ttn = 55) while showing comparable productivity (Table 39.3, entry 10) [46]. Current efficiency was also slightly improved. Once again, high enantiomeric excesses were reported (ee = 98%). A third approach is the use of an aqueous-organic O

Cathode

[Cp*Rh(bpy)H2O] 2+

NADPH

H+ S-ADH OH

H2O

H2O [Cp*Rh(bpy)H]+

NADP+

FIgURE 39.6 Asymmetric reduction of a prochiral ketone with an ADH and Cp*Rh(bpy) as mediator.

© 2016 by Taylor & Francis Group, LLC

1524

Organic Electrochemistry

two-phase system (Table 39.3, entry 11) [46]. In a reaction setup similar to the one described earlier, methyl tert-butyl ether was added as the organic phase, which served as reservoir for the substrate and for in situ product extraction, resulting in higher turnover numbers (ttn = 64). A product solution of 180 mM (R)-phenylethanol with an ee > 99.9% was obtained, thus clearly indicating the potential of combining the concept of two-phase systems with electroenzymatic synthesis. It is worth mentioning that among the mediators developed Cp*(Rh(bpy) is very privileged, as it can reduce both the phosphorylated (NADP+) and unphosphorylated forms (NAD+) of the nicotinamide cofactor. Apart from electrochemical activation, it can also use other chemical redox equivalents (hydrogen, formate, phosphite) [14,47,48]. The electrochemical activity can be tuned by introducing substituents on the bipyridine moiety. A recent screening found complexes with threefold higher activity for NADP+ reduction than the unmodified mediator [49]. The best ligands in terms of activity were the 5,5′-dimethyl- and 4,4′-methoxy-bipyridines. However, these complexes were not tested in enzyme reactions. The Cp*Rh(bpy) mediator also has the severe drawback that the enzyme in its presence is instable. It was speculated that this instability of enzymes in the presence is due to basic amino acids (lysine, histidine, and arginine) present in the protein [50]. The side chains of these amino acids might interfere with the coordination sphere of the rhodium and lead to enzyme aggregates or inactivation. Moreover, it was known that enzymes with sulfur-containing amino acids (cysteine and methionine) can react with oxygen and heavy metals, both leading to destabilization. An electroanalytical investigation of the mediator in the presence of all proteinogenic amino acid revealed a strong irreversible inactivation in the presence of cysteine (which was to be expected) and an equilibrium-based (dependent on the concentration of the amino acid) inactivation in the presence of histidine and tryptophan, which was not previously known [51]. The impact of different amino acids on the mediator is shown in Figure 39.7. For details on the electroanalytical investigations, the reader is referred to the supporting information of Hildebrand and Lütz [51]. This was the first comprehensive study on the mutual inactivation phenomenon taking all amino acids into account. One could speculate that the N-heterocyclic nature of the amino acid side chain rather than its nucleophilicity dominate its interaction with the mediator. This has direct implications for the reaction setup: buffer or pH change will probably not alleviate the problem. Peak current in presence of amino acid

1.00 0.90 0.80

HS 0.70

OH NH2 cysteine

0.60 Ip/Ipo/-

O OH

Peak current in absence of amino acid O N O N H

NH2

OH

N

NH2

tryptophan

histidine

0.50 0.40 0.30 0.20 0.10

FIgURE 39.7 Impact of different amino acids on the mediator activity.

© 2016 by Taylor & Francis Group, LLC

o Se r Th r Tr p Ty r Va l

Pr

u Ly s M et Ph e

Le

A

la A rg A sn A sp Cy s G lu G ln G ly H is lle

0.00

1525

Electroenzymatic Synthesis

4

1. Pump 2. Membrane filtration 3. Electrochemical cell 4. Immobilized ADH polymer-enlarged mediator

H 2O Rh N

N

2 3

1

Cathode

OH Polymer-Rh(bpy)red

NADP+ CI O

ADH Polymer-Rh(bpy)ox

NADPH CI

FIgURE 39.8 Reaction and reactor setup with a spatially separated electroenzymatic reduction with polymer-enlarged mediator.

The mutual inactivation can only be avoided if mediator and enzyme are fully separated. This was achieved using a polymer-enlarged mediator retained in an electrochemical membrane reactor together with an immobilized enzyme (Table 39.3, entry 12; see also Figure 39.8). This is, to the best of our knowledge, the first example of an electrochemical activation of the Cp*Rh(bpy) mediator coupled to an enzymatic reaction, where 100% of the enzyme activity could be recovered. The ttn for the mediator was 200, which is the highest value for the Cp*Rh(bpy) system under electrochemical activation (with formate, higher ttns have been obtained). An approach that was thought to alleviate the stability problems is the use of nitrogen containing buffers. Höllrigl et  al. used the Cp*Rh(bpy) mediator together with an ADH from Thermus sp. for the production of (1S,3S)-3-methylcyclohexanol at a diastereomeric excess of 96% from the corresponding racemic ketone (Table 39.3, entry 13) [52]. In this case, octane was used as a second organic phase. However, the stabilizing effect of nitrogen-containing buffers could not be confirmed with other enzymes [51]. As it is very difficult to find an electrochemical redox catalyst that fulfils all requirements for regenerating NAD(P)H effectively, attempts have been made to regenerate the cofactor indirectly by coupling with a second enzymatic reaction in addition to an electrochemical reaction step. While this overcomes the problem of dimer formation, it loses the potential advantage associated with electroenzymatic synthesis over alternative purely enzymatic cofactor regeneration approaches, which also use a second enzyme but do not need the electrochemical equipment (see Table 39.1). Several reactions are reported where methyl viologen (MV2+) in combination with diaphorase, lipoamide dehydrogenase (LipDH), ferredoxin NADP+ reductase (FNR), or so-called AMAPORS acts as electron shuttle to regenerate NAD(P)H and will be summarized in the following paragraphs. Reduction of NADH with MV2+ together with diaphorase was used in the ADH catalyzed reduction of cyclohexanone (Table 39.3, entry 14) [53], 2-methylcyclohexanone (Table  39.3,  entry 15)  [53],

© 2016 by Taylor & Francis Group, LLC

1526

Organic Electrochemistry

and 3-methylcyclohexanone (Table 39.3, entry 16) [53]. MV2+, diaphorase, NAD+, and the ADH enzyme were immobilized on the electrode. A potential of −0.8 V versus SCE was applied for electrolysis. A yield of >99% and a current efficiency of 97.6% were obtained for the synthesis of cyclohexanol. For the synthesis of (1S,2S)-(+)-2-methylcyclohexanol, a yield of 48.9% was achieved with an ee of >99%, a turnover of 91 for MV2+ and a current efficiency of 98.6%. For (1S,3S)-(−)-2methylcyclohexanol, a yield of 51.7% with an ee of >93.1%, a turnover of 94 for MV2+ and a current efficiency of 96.5% was obtained. The regeneration system MV2+/diaphorase is also used in combination with NAD+ and d-lactate dehydrogenase (Table 39.3, entry 17) [54]. In this reaction, pyruvate is reduced to d-lactate. By adding 1.5 U mL−1 diaphorase, 0.2 mM MV2+ , and 0.3 mM NAD+ as well as 50 mM pyruvate to the solution and performing an electrolysis at −0.7 V versus Ag|AgCl, a conversion of roughly 80% was achieved. The synthesis of (R)-mandelic acid from benzoylformate catalyzed by benzoylformate dehydrogenase was carried out with regeneration by MV2+ and diaphorase (Table 39.3, entry 18) [55] as well as with FAD and lipoamide dehydrogenase (LipDH, Table 39.3, entry 19) [55]. Regeneration with MV2+ and diaphorase led to a conversion of 80% in 30 h and the total turnover number for NAD+ was 133. The FAD-LipDH system provided much better results; with this regenerating system a conversion of 95% in 18 h and a turnover number of 158 for the cofactor was obtained. MV2+ was used together with LipDH for the synthesis of d-lactate by d-lactate dehydrogenase (Table 39.3, entry 20) [56]. At a tungsten electrode, a potential of −0.72 V versus SCE was applied. In a batch experiment, a conversion of 81% was obtained with an ee of 94% for the product and a turnover number of 940 for NAD. The combination of MV2+/LipDH has also been tested in a continuous process. Bergel and coworkers applied this in their D-MER together with an ADH for the synthesis of cyclohexanol from cyclohexanone (Table 39.3, entry 21) [45]. Electrolysis was carried out at −0.7 V versus SCE. During this synthesis, the enzyme reaction between NAD+ and the mediator seemed to be the ratelimiting step as the conversion could be increased from 26% to 65% with the further addition of LipDH. Together with ferredoxin reductase (FDR), MV2+ was used for the synthesis of d-lactate from pyruvate by d-lactate dehydrogenase (Table 39.3, entry 22) [56]. Turnover numbers of 900 for NADH, 2.2 ∙ 107 for d-lactate dehydrogenase and 7.3 ∙ 106 for FNR were achieved. The application of the same regeneration system for the production of l-glutamate from 2-oxoglutarate by l-glutamate dehydrogenase (Table 39.3, entry 23) [56] led to turnover numbers of 1000 for NADPH, 1.1∙107 for l-glutamate dehydrogenase and 7.5∙106 for FNR with complete conversion of the substrate. AMAPORs were used in combination with both MV2+ (Table 39.3, entries 24 and 25) [32] and cobalt sepulchrate (CoSep) for the synthesis of (S)-glutamate (Table 39.3, entry 26) [32]. Dissolved in TRIS buffer (pH 7.3) and electrolyzed at a potential of −729 mV versus SCE were 100  mM 2-oxoglutarate, 250 mM ammonium acetate, 3 mM MV2+, and 36 U of glutamate dehydrogenase. Produced with 0.5 mM NADP+ and 50 mg wet cells of C. thermoaceticum was 29 mol product kgcatalyst−1 h−1 was, whereas for 0.5  mM NAD+ and 100 mg of wet cells, a productivity of 7 mol product kgcatalyst−1 h−1 was obtained. Conversion varied between 95% and 99%. When using CoSep instead of MV2+, the performance dropped to only 3.5% of the productivity reached before, which indicates that the use of MV2+ is beneficial (Figure 39.9). Acetophenone can be reduced directly to the racemic mixture of the corresponding alcohol at a potential of ~−0.8 V versus Ag|AgCl. The corresponding alcohol can be used as a substrate for indirect enzyme-coupled NADPH regeneration without adding a second enzyme. This approach has been tested in the synthesis of different alcohols leading to ee values of >99%, but only with a conversion of around 10% (Table 39.3, entry 27) [57]. The same substrates have also been used for synthesis reactions with MV2+ and FDR or diaphorase (Table 39.3, entry 28) [57]. Conversion was doubled by these regeneration systems and the current efficiencies were increased.

© 2016 by Taylor & Francis Group, LLC

Electroenzymatic Synthesis

NH2

O +

R–N

+

N –R

NH2

HOOC

Cathode

N R

H2O

AMAPOR

GluDH

+

+

NH2

N–R

O

NH3

O R–N

COOH

HOOC

COOH

N R

FIgURE 39.9

Reductive amination of 2-oxoglutarate using an AMAPOR and glutamate dehydrogenase (GluDH).

1527

© 2016 by Taylor & Francis Group, LLC

1528

Organic Electrochemistry

A hydrogenase from Alcaligenes eutrophus was used to regenerate NADH in the reduction of 2-oxoglutarate to l-glutamate by l-glutamate dehydrogenase (Table 39.3, entry 28) [58]. In a 5 mL cylindrical divided cell, electrolysis was carried out at a potential of −0.7 V versus SCE. The working and auxiliary electrodes were made of platinum. Complete conversion of 300 mM substrate was achieved with a turnover frequency (tof) of approximately 200 h−1.

IV. FLAVIN-DEPENDENT OXIDATIONS WITH OXIDATIVE REgENERATION In contrast to the nicotinamide cofactors discussed in Sections II and III, FAD is typically bound to the enzyme more strongly and cannot diffuse freely to the electrode. In oxidase enzymes, the cofactor can be easily regenerated with molecular oxygen. Typically hydrogen peroxide is formed as a by-product, which may decrease enzyme activity and stability [59]. An approach that can be used to overcome this problem is the addition of catalase to remove the hydrogen peroxide [60,61]. Electrochemistry provides an elegant anaerobic method for cofactor regeneration without using catalase. The large size of the biomolecules and, in most cases, the location of the active site, where the FAD is bound, deep inside the protein layer hampers direct electron transfer [62]. Thus, for an effective electron transfer from the FAD to the electrode, it is necessary to use mediators as electron shuttles. Ferrocene derivatives are most often used for this purpose. For example, ferroceneboronic acid was used in the production of p-hydroxybenzaldehyde by oxidation of p-methylphenol with p-cresolmethyl hydroxylase (PCMH) (Table 39.4, entry 1) [63]. The reaction was carried out in a two-compartment cell with a gold working electrode. A yield of up to 85% was obtained. A similar electroenzymatic reaction with polymer-enlarged ferrocene derivative and a graphite foil as electrode yielded 84% p-hydroxybenzaldehyde (Table 39.4, entry 2) [64]. Electrochemical regeneration of FAD was also carried out in a continuous process [19,65,66]. An electrochemical cell with a carbon felt working electrode was connected to an enzyme membrane reactor forming a so-called EEMR. The product solution is passed through an ultrafiltration membrane, where the enzyme is held back in the flow cycle. As the mediator size is similar to that of the educts and products, it has to be bound to a homogeneously soluble polymer to ensure that it remains in the reaction cycle. With this system, the oxidation of p-methylphenol to p-hydroxybenzaldehyde was performed with p-cresolmethyl hydroxylase (PCMH) as biocatalyst (Table 39.4, entry 3, see Figure 39.10) [65]. The reaction was carried out in a galvanostatic mode and after 50  h a steady state with almost quantitative conversion of p-methylphenol was reached. Ttns of 400,000 for the enzyme and 500 for the mediator were achieved. In the same reactor, the substrate 4-ethylphenol was converted by PCMH, leading to (S)-1-(4-hydroxyphenyl)-ethanol with 88% ee (Table 39.4, entry 4) [65]. Hydroxyacetophenone from the oxidation of benzylic alcohol was found as a by-product. The process was also carried out with the enzyme 4-ethylphenol methylenehydroxylase (EPMH) (Table 39.4, entry 5) [65]. The (R)-enantiomer was obtained in 99% ee and the ketone by-product formation decreased to 10%. Galactose oxidase was used for the oxidation of xylitol to l-xylose with regeneration by ferrocene (Table 39.4, entry 6) [66]. Fast enzyme deactivation was observed in batch electrolyses due to denaturation at the counter electrode and the influence of shear forces. Therefore, a fixed-bed reactor with immobilized galactose oxidase was coupled with an electrochemical flow-through cell. The substrate concentration was 7.4 mM in this case and the conversion was completed within 3 weeks leading to a ttn of > 200,000 for the enzyme. l-Glycerol-3-phosphate oxidase (GPO) was used in the batch electrolysis of l-glycerol-3-phosphate forming dihydroxyacetone phosphate (DHAP) (Table 39.4, entry 7, see Figure 39.11) [19]. Again, polymer-bound ferrocene served as the mediator. The DHAP formed was used in situ for an aldolase-catalyzed C–C bond formation reaction. The enzyme used for the second step was fructose-1,6-diphosphate aldolase.

© 2016 by Taylor & Francis Group, LLC

Entry 1

Enzyme

Substrate

PCMH from Pseudomonas alcaligenes u.a. (PCMH, E.C. 1.17.99.1) PCMH from Pseudomonas putida (PCMH, E.C. 1.17.99.1)

p-methylphenol

3

PCMH from Pseudomonas putida (PCMH, E.C. 1.17.99.1)

4

Product

Mediator

Experimental Data

Literature

p-hydroxy benzaldehyde p-hydroxy benzaldehyde

azurine, ferrocene boric acid

Conversion 85%

[63]

PEG-ferrocene

[64]

p-methylphenol

p-hydroxy benzaldehyde

PCMH from Pseudomonas putida (PCMH, E.C. 1.17.99.1)

p-ethylphenol

Conversion ~70%, ee 88%

[65]

5

EPMH from Pseudomonas putida EPMH

p-ethylphenol

(S)-1-(4hydroxyphenyl)ethanol (R)-1-(4-hydroxy phenyl)ethanol

PEG-ferrocene, α,ω,bismethylferrocene, polyethyleneglycol (20,000) PEG-ferrocene

Conversion 84% in 17 h, ttnenzyme 13,0000, ttnmediator 66, current efficiency 100% Conversion 100%, ttnenzyme 400,000, ttnmediator 500

galactose oxidase from Fungus fusarium NRRL 2903 (GOase, E.C. 1.1.3.9)

l-xylitol

l-xylose

7

l-glycerol-3-phosphate oxidase from Pediococcus sp. (E.C. 1.1.3.21) d-amino acid oxidase (dAAO) from Trigonopsis variabilis Crude enzyme extract from Escherichia coli Crude enzyme extract from Saccharomyces cerevisiae

l-glycerol-3phosphate dl-methionine

dihydroxyacetone phosphate l-methionine

Conversion 100%, ee 93% (optical purity) and 99% (GC), respectively 7.4 mmol L−1 in 3 weeks, ttnenzyme 208,720, ttnmediator 15, current efficiency 34.6% 90 mM

[65]

6

PEG-ferrocene, α,ω,bismethylferrocene, polyethylene glycol (20,000) PEG-ferrocene, α,ω,bismethylferrocene, polyethylene glycol (20,000) PEG-ferrocene

fumarate xylose

Productivity 27 g L−1 d−1, ee > 99.9%, tofFcCOOH 12.9 h−1 ~0.8 mM in 6 h ~0.55 mM in 10 h

[67]

succinate xylitol

ferrocene carboxylic acid (FcCOOH) None None

2

8 9 10

p-methylphenol

Electroenzymatic Synthesis

TAbLE 39.4 Syntheses with Oxidative Regeneration of Flavin-Dependent Enzymes

[65]

[66]

[19]

[70] [70]

1529

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Organic Electrochemistry O

H

2 PEG-Fc +

PCMH-FADH2

OH

PCMH-FAD

2 PEG-Fc

PCMH-FADH2

2 PEG-Fc +

PCMH-FAD

2 PEG-Fc

Anode

OH

OH

CH3

H2O

OH

FIgURE 39.10 Oxidation of p-methylphenol using FAD dependent PCMH and a polymer-bound ferrocene mediator (PEG-Fc) for cofactor regeneration. H

OH CH2N(CH3)3+

FAD

Fe

FADH2

Fe+

Anode

OPO32–

HO

GPO O OPO32–

HO

CH2N(CH3)3+

OH O Aldolase

H OH

O

OH

O

HO

Phosphatase HO HO

OH

OPO32–

HO OH HO

OH

FIgURE 39.11 Synthesis of dihydroxyacetone phosphate (DHAP) with glycerol-3-phosphate oxidase (GPO) and ferrocene as mediator and its subsequent use in an aldolase-catalyzed C–C bond formation reaction.

Ferrocene carboxylic acid is more water-soluble than unmodified ferrocene and has been used in the resolution of dl-methionine catalyzed by a d-amino acid oxidase (DAAO) from T. variabilis (Table 39.4, entry 8) [67]. l-Methionine was obtained from racemic methionine with an optical purity of ee > 99.9%. By adding 10 vol% of the ionic liquid 1,3-dimethylimidazolium dimethylphosphate [MMIM][Me2PO4] the productivity could be increased 1.5-fold from 18 g L−1 d−1 in aqueous buffer solution to 27 g L−1 d−1 in the presence of the ionic liquid. This is, together with

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a peroxidase-catalyzed reaction reported in the same publication, the first example of using ionic liquids as additives in electroenzymatic synthesis. Ionic liquids, salts that are liquid at ambient temperatures, have found widespread use as cosolvents in biocatalysis (e.g., [68,69]). Depending on their chemical structure, they can both serve as a supporting electrolyte and solubilizing agent for poorly water-soluble substrates. An example of an electrochemical regeneration of FAD without a mediator is the use of a catalytic electrode for the oxidation of succinate to fumarate (Table 39.4, entry 9) [70] and xylitol to xylose (Table 39.4, entry 10) [70] with crude enzyme extract from E. coli and S. cerevisiae, respectively. Graphite powder, inorganic binder and 3% Cu(II) ion were mixed and fixed on a plate. The counter electrode was prepared in the same way with Fe(II) ions instead of copper ions. The reactions were carried out in a two-compartment cell with 2 V cell voltage.

V. FLAVIN-DEPENDENT REACTIONS WITH REDUCTIVE REgENERATION Monooxygenases catalyze a broad range of oxidation reactions, for example, hydroxylations, epoxidations, heteroatom oxidations, and Baeyer–Villiger oxidations [71]. The oxidation proceeds via transfer of an oxygen atom to the substrate, so that molecular oxygen is essential as cosubstrate. The second atom of molecular oxygen is reduced to water. Therefore, monooxygenases need both an aerobic environment and reductive cofactor regeneration (either FADH2 directly or via NAD(P)H). Monooxygenases can be subdivided into several classes according to their cofactors and subdomain structure [72,73]. A couple of electroenzymatic syntheses dealing with these enzymes have been reported. For the electroenzymatic ω-hydroxylation of lauric acid (Table 39.5, entries 1 and 2) [74,75], a fusion protein of a cytochrome P450 monooxygenase and rat NADPH-P450 reductase as well as a reconstitution system with an engineered monooxygenase and added NADPH-P450 reductase were used. Two electrons are required to activate P450 for monooxygenation, which are typically derived from NADPH. The fusion proteins are able to accumulate two or more electrons donated by an electrochemical mediator, in this case cobalt(III)sepulchrate trichloride [74]. 1,1′-Dicarboxycobaltocene was used in a similar way for the monooxygenation of lauric acid with cytochrome P450 BM3 [75]. This worked both with the holoprotein and the isolated heme; however, the reaction with the heme domain was not very efficient due to overoxidation. 2-Hydroxybiphenyl-3-monooxygenase catalyzes the specific o-hydroxylation of α-substituted phenol derivatives. An approach with indirect electrochemical regeneration of NADH used a [Cp*Rh(bpy)Cl]Cl complex as mediator (Table 39.5, entry 3) [76]. This system is quite complex and requires several components (see Figure 39.12). An undesired side reaction was the formation of hydrogen peroxide by the reaction of the hydridorhodium complex with molecular oxygen and reduction of dissolved oxygen at the carbon-felt cathode. The productivity of the electroenzymatic reaction was 204 mg L−1 h−1. Direct unmediated electron transfer to FAD without inclusion of the NADH regeneration cycle was performed in the epoxidation of styrenes catalyzed by styrene monooxygenase (see Figure 39.13) [77]. Several (S)-epoxidized styrenes were obtained with ee > 98% at rates between 0.074 and 0.222 mM h−1 (Table 39.5, entry 4). Both for the hbpA and styA reaction, comparison data exist for either whole-cell, enzymatic, or chemical regeneration approaches. The electrochemical regeneration typically was inferior to the alternative approaches either due to enzyme instability, reaction velocity, or ttns. The natural function of the pyruvate dehydrogenase complex (PDC) is to decarboxylate pyruvate to acetyl-coenzyme A. The reverse reaction, formation of pyruvate, was performed by the indirect electroenzymatic fixation of CO2 in acetyl-coenzyme A with PDC (Table 39.5, entry 5) [78] using MV2+ as mediator. Pyruvate production stagnated at a certain time due to the formation of coenzyme A, which inhibited the reaction. This inhibition was circumvented by adding phosphotransacetylase and acetyl phosphate, which converted the inhibitory coenzyme A into

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TAbLE 39.5 Syntheses with Reductive Regeneration of Flavin-Dependent Enzymes Process 1 2 3

4

5 6

l-amino acid oxidase from Crotalus adamanteus venom (l-AAO, E.C. 1.4.3.2)

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Substrate

Product

Cofactor

lauric acid

ω-hydroxy lauric acid

FAD+/H2

lauric acid

ω-hydroxy lauric acid

FAD+/H2

2-hydroxybiphenyl

2,3-dihydroxybiphenyl

FAD+/H2, NADP+/H

styrene and derivatives

(S)-epoxidized styrene derivatives

acetyl coenzyme A pyruvic acid

phenylpyruvic acid

Mediator

Experimental Data

Literature

cobalt(III)sepulchrate trichloride 1,1′-dicarboxycobaltocene [Cp*Rh(bpy)Cl]Cl

0.65 nmol min−1 mL−1

[74]

ttnenzyme 224

[75]

204 mg h−1 L−1, tofmediator 11 h−1

[76]

FAD+/H2

None

0.074–0.222 mMh −1, ee 98.1–99.9%

[77]

pyruvate d-alanine

FAD FAD+/H2

FAD+/H2

0.98 µmol, in 100 h 8.9 mM in 10 h, ee > 99%, ttnmediator 36,000, current efficiency ~ 97% 8.5 mM in 10 h, ee > 99%

[78] [79]

l-phenylalanine

MV2+ 1-aminopropyl-1′methyl-4,4′dipyridinium iodide 1-aminopropyl-1′methyl-4,4′dipyridinium iodide

[79]

Organic Electrochemistry

7

Enzyme fusion protein rFP450 [mRat4A1/ mRatOR]L1 (rFP4504A1) cytochrome P450 BM3 from Bacillus megaterium (BM3) 2-hydroxybiphenyl-3monooxygenase from Pseudomonas azelaica (HbpA) styrene monooxygenase (only FADH2-dependent oxygenase component) from Pseudomonas sp. VLB 120 (StyA) PDC d-amino acid oxidase from porcine kidney (d-AAO, E.C. 1.4.3.2)

1533

Electroenzymatic Synthesis OH [Cp*Rh(bpy)H2O]2+

NADPH

H+ Cathode

O2

hbpA HO H2O

H2O

OH

H2O

NADP+

[Cp*Rh(bpy)H]+

FIgURE 39.12 Electroenzymatic aromatic hydroxylation catalyzed by 2-hydroxybiphenyl-3-monooxygenase (hbpA).

O

FAD Cathode

H2O styA

O2 FADH2

FIgURE 39.13 Electroenzymatic epoxidation catalyzed by styrene monooxygenase (styA) with direct flavin (FAD) reduction.

acetyl-coenzyme A. Thus, the pyruvate production was five times higher, but a maximal production of 0.98 µmol pyruvate was achieved within 100 h. The asymmetric synthesis of amino acids was performed with enzyme and mediator immobilized on a glassy carbon electrode (Table 39.5, entries 6 and 7) [79]. The FAD-dependent enzyme amino acid oxidase (AAO) oxidizes amino acids to imino acids, which consecutively (spontaneously) hydrolyze to the corresponding α-keto acids. The reaction was reversed by using 1-aminopropyl1′-methyl-4,4′-dipyridinium iodide as mediator coimmobilized on the electrode. If a potential is applied at which the mediator is permanently in its reduced state, the FAD will be provided in its reduced form. This converts the AAO into a reducing agent for imino acids. With this system, d-alanine and l-phenylalanine were synthesized from pyruvic acid and phenylpyruvic acid, respectively, with >99% ee and yields of around 30% within 10 h.

VI.

ELECTROENZyMATIC OXIDATIONS WITH PEROXIDASES

Peroxidases are not dependent on NAD(P) or FAD cofactors, but contain other prosthetic groups where the redox reaction takes place, including vanadate or heme (see Figure 39.1). They require a cosubstrate—generally hydrogen peroxide—to bring about the reaction and regenerate the prosthetic group to the desired redox state. High levels of the oxidant are however detrimental to enzyme stability in some cases. Especially the porphyrin unit in heme peroxidases is very sensitive

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Organic Electrochemistry

to oxidative degradation by hydrogen peroxide. This can be circumvented by reaction engineering approaches, for example, sensor-controlled dosing of H2O2 [80], but depending on the quality of the sensor, it is impossible to avoid high local concentrations of the oxidant and the addition of diluted hydrogen peroxide continuously increases the reaction volume. Alternatively, the hydrogen peroxide can be synthesized in situ by a chemical reductant or a second enzyme [81]. The general drawbacks of these methods are the formation of a coproduct in stoichiometric amounts and the difficulty of adjusting the rate of hydrogen peroxide generation for every single enzymatic reaction to avoid accumulation of the cosubstrates. The electrochemical reactant supply is a useful method for the controlled addition of hydrogen peroxide. Cathode materials that can be used for the generation of the oxidant via the reduction of dissolved oxygen are mercury, gold, and carbon [82] since with most other materials the formation of water predominates. Due to safety risks while working with mercury and the high cost of gold electrodes, carbon is the favored electrode material. The electrochemical hydrogen peroxide supply for the enzyme avoids the formation of by-products and does not increase the reaction volume. A further advantage is that the rate of hydrogen peroxide formation can be adjusted by simply varying the potential. The combination of H2O2 generation with a peroxidase was for instance used in the conversion of 2,4,6-trimethylphenol to 3,5-dimethyl-4-hydroxybenzaldehyde by horseradish peroxidase (HRP) in a divided cell batch reactor and a flow cell reactor (Table 39.6, entries 1 and 2) [83]. The regioselectivity of the oxidation depended on the applied potential; nevertheless a mixture of at least four different products was always formed. The same enzyme was also used for the dealkylation of N,N-dimethylaniline to N-methylaniline (Table 39.6, entry 3) [84]. The oxidation of veratryl alcohol to veratraldehyde is described using lignin peroxidase with in situ electrogeneration of hydrogen peroxide (Table 39.6, entry 4) [85]. In this reaction, the anode (a platinum plate) and cathode (consisting of reticulated vitreous carbon) were divided by a cation exchange membrane and the cathode compartment was continuously aerated with oxygen. The most frequently used peroxidase in organic synthesis is chloroperoxidase (CPO). This heme peroxidase catalyzes a broad range of oxidation reactions [86]. In a first attempt, the chlorination of barbituric acid to 5-chlorobarbituric acid was investigated (Table 39.6, entry 5) [87]. The reactor consisted of an electrochemical cell in which oxidation of H2O to O2 is followed by the generation of H2O2 and a hollow-fiber filtration module, in which the biocatalysis takes place. The product was retained by an ion exchange resin. After complete conversion, 96% of the product was recovered from the column. As CPO is known to oxidize various sulfides enantioselectively, in situ generation of hydrogen peroxide was applied to the synthesis of various sulfoxides, for example, (R)-methylphenylsulfoxide (Table 39.6, entry 6) [88], (R)-methyl p-tolylsulfoxide (Table 39.6, entry 7) [89], (R)-methoxyphenyl methyl sulfoxide (Table 39.6, entry 8) [89], and N-MOC(methoxycarbonyl)-l-methionine methyl ester sulfoxide (Table 39.6, entry 9) [89]. Best results were achieved for the synthesis of (R)-methylphenylsulfoxide. A space–time yield of 30 g L−1 d−1 was obtained in a divided batch cell using a carbon felt cathode fixed on stainless steel and an anode compartment formed by a platinum wire in a dialysis sack. The product had an ee of 98.5% (Figure 39.14). Using a 3D electrolysis cell with a bed of graphite beads as cathode, a productivity of 104 g L −1 d−1 of the chiral sulfoxide was achieved (Table 39.6, entry 10) [90]. Also in this case, the use of ionic liquids has a synergistic effect (Table 39.6, entry 11) [67], as the sulfides used as substrates usually have low aqueous solubilities and small amounts of ionic liquids increased productivity from 18 g L−1 d−1 (in plain aqueous buffer without cosolvent) to a maximum of 75 g L −1 d−1 in the presence of 2 vol% of the ionic liquid 1-ethyl-3-methylimidazolium ethylsulfate [EMIM][EtSO 4].

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Entry

Enzyme

Substrate

Product

1

horse radish peroxidase (HRP, E.C. 1.11.1.7)

2,4,6-trimethylphenol

Mixture of substances

2

horse radish peroxidase (HRP, E.C. 1.11.1.7)

2,4,6-trimethylphenol

Mixture of substances

3 4 5

Horse radish peroxidase (HRP, E.C. 1.11.1.7) lignin peroxidase from P. chrysosporium (LiP) CPO from Caldariomyces fumago (CPO, E.C. 1.11.1.10) CPO from Caldariomyces fumago (CPO, E.C. 1.11.1.10) CPO from Caldariomyces fumago (CPO, E.C. 1.11.1.10) CPO from Caldariomyces fumago (CPO, E.C. 1.11.1.10) CPO from Caldariomyces fumago (CPO, E.C. 1.11.1.10) CPO from Caldariomyces fumago (CPO, E.C. 1.11.1.10) CPO from Caldariomyces fumago (CPO, E.C. 1.11.1.10)

N,N-dimethylaniline veratryl alcohol barbituric acid

N-methylaniline veratraldehyde 5-chlorobarbituric acid

thioanisole

(R)-methylphenyl sulfoxide (R)-methyl p-tolyl sulfoxide (R)-methoxyphenyl methyl sulfoxide N-MOC-l-methionine methyl ester sulfoxide (R)-methylphenyl sulfoxide (R)-methylphenyl sulfoxide

6 7 8 9 10 11

methyl p-tolyl sulfide 1-methoxy-4(methylthio)benzene N-MOC-l-methionine methyl ester thioanisole thioanisole

Experimental Data

Literature

40% of 3,5-dimethyl-4-hydroxybenzyl alcohol, traces of 3,5-dimethyl-4hydroxybenzaldehyde and 2,6-dimethylbenzoquinone 68% of 3,5-dimethyl-4-hydroxybenzyl alcohol, 3% 3,5-dimethyl-4-hydroxybenzaldehyde, 3% 4-(4-hydroxy-3,5-dimethyl-benzyloxy)-2,4,6trimethylcyclohex-2,5-dienone Conversion 90% Not available Conversion > 96%, ttnenzyme 500,000, current efficiency > 90% 30 g L−1 d−1, ee 98.5%, ttnenzyme 95,000, current efficiency 65.6% Conversion 76%, ee 93%, ttnenzyme 58,900

[83]

[89]

Conversion 83%, ee 99%, ttnenzyme 64,400

[89]

Conversion 60%, dr 81:19 (R), ttnenzyme 700

[89]

104 g L−1 d−1, ee 98.5%, ttnCPO 145,000

[90]

75 g L−1 d−1, ttnCPO 123,000

[67]

Electroenzymatic Synthesis

TAbLE 39.6 Synthesis Reactions with Electrochemical Cosubstrates Supply

[83]

[84] [85] [87] [88]

1535

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Organic Electrochemistry O S O2

Cathode

H2O CPO

S H2O2

FIgURE 39.14

VII.

Electroenzymatic synthesis of chiral sulfoxides catalyzed by CPO.

SyNTHESIS REACTIONS WITH ELECTROCHEMICAL SUbSTRATES SUPPLy OR PRODUCT CONVERSION

In contrast to the reactions reported in the previous sections, electrochemistry can not only be used for cofactor regeneration, but also to establish reaction cascades where the substrate of an enzymatic reaction is generated electrochemically or the product of the biocatalytic step serves as the educt for a consecutive electrochemical reaction. In some cases, this allows to establish paired electrolyses, where both electrodes serve as productive working electrodes [20]. The approach of combining an enzymatic reaction with paired electrolysis was performed in the synthesis of malate from lactate (Table 39.7, entry 1) [91]. In the anodic half-cell reaction, lactate is oxidized to pyruvate using an RuO2/Ti electrode. The formed pyruvate migrates through TAbLE 39.7 Syntheses with Electrochemical Substrates Supply/Product Conversion Entry 1

2

3

4

Enzyme malic enzyme (ME, E.C. 1.1.1.40) l-lactate dehydrogenase (l-LDH, E.C. 1.1.1.27) l-alanine dehydrogenase from Bacillus subtilis (l-Ala DH, E.C. 1.4.1.1) d-AAO from T. variabilis

© 2016 by Taylor & Francis Group, LLC

Substrate

Experimental Data

Literature

MV

~8 µmol in 20 h

[91]

NAD+/H

Direct

[92]

d-alanine

NAD+/H

Direct

Conversion ~97%, 15 mmol L−1 d−1, ttncofactor 200, tofcofactor 6.3 h−1 10 mM in 140 h

[93]

l-leucine

FAD

N.A.

3.5 mmol L−1 d−1

[94]

Product

Cofactor

Mediator

lactate (pyruvate)

malate

NADP+/H

l-lactate

d-lactate

l-alanine

d-leucine, dl-leucine or 4-methyl-2oxovaleric acid

1537

Electroenzymatic Synthesis

an anion exchange membrane into the cathode compartment, where the enzymatic reaction takes place. Malic enzyme (ME) converts the produced pyruvate to malate under addition of CO2. The cofactor NADPH is regenerated electrochemically at the cathode with MV2+ as mediator and ferredoxin-NADP+-reductase as redox partner. Around 8 µmol of malate was produced in 20 h with this setup, the rate limiting steps being the enzymatic reaction as well as mass transfer through the membrane. Another combination of electrochemical cofactor regeneration and electrogeneration of the substrate is the conversion of l-lactate into d-lactate by an l-lactate dehydrogenase (l-LDH) catalyzed oxidation of the l-antipode into pyruvate (see Figure 39.15). This oxidation is combined with the reduction of the pyruvate at a mercury cathode, forming racemic d,l-lactate (Table 39.7, entry 2) [92]. From the racemic mixture produced in this cathodic half reaction, the l-enantiomer undergoes the enzymatic reaction and a continuous oxidation–reduction cycle proceeds. The anode was used for NAD+ regeneration. The anodic and the cathodic reaction were carried out in two separate electrochemical cells. It was possible to achieve a conversion of 97% with this setup. In a similar setup, d-alanine was produced from l-alanine with l-amino acid dehydrogenase (AADH) (Table 39.7, entry 3) [93]. The imino acid easily hydrolyzes to the α-keto acid, since the equilibrium of this reaction highly favors the keto acid over the imino acid in aqueous solution. Thus, the use of a mercury cathode was necessary for the selective reduction of the imino acid and to avoid formation of the α-hydroxy acid via reduction of the keto compound. Altogether the electrochemical conversion to the amino acid was very slow. Due to this fact, the overall process needed a long time, in fact after 20 days, only 60% of d-alanine was produced. Recently, the electrochemical reduction of the imino acid was combined with an d-amino acid oxidase (d-AAO) rather than a dehydrogenase enzyme (Table 39.7, entry 4) [94]. This is favorable as the oxidase requires only aeration with molecular oxygen for cofactor regeneration instead of nicotinamide regeneration in the case of the dehydrogenase, and a paired electrolysis is not necessary (see Figure 39.16). As a consequence, the productivity obtained in this setup is one order of magnitude higher (3.5 mmol L −1 d−1) than with the dehydrogenase enzyme. Moreover, this reaction cycle can be started from the α-keto acid, the racemic amino acid, or the d-amino acid and therefore is a strategy for asymmetric synthesis, deracemization, or stereoinversion, respectively. O NAD+

Cathode

Anode

COOH

- LDH

OH NADH COOH OH

COOH

FIgURE 39.15 Electroenzymatic synthesis of d-lactic acid in a paired electrolysis with cofactor regeneration and pyruvate reduction.

© 2016 by Taylor & Francis Group, LLC

1538

Organic Electrochemistry

Cathode

–H2O +NH3

O R

COOH

+H2O –NH3

NH

NH2

NH2

+ R

COOH

R

COOH

R

COOH

D - AAO

FIgURE 39.16

VIII.

Electroenzymatic synthesis of amino acids using an d-AAO.

SUMMARy AND OUTLOOK

In the previous sections, the current state of the art in the field of research of electroenzymatic syntheses is presented. The vast majority of the processes use electrochemistry for the regeneration of the cofactors of the corresponding oxidoreductase enzymes. In nearly all cases, indirect electron transfer between electrode and cofactor was applied due to the very slow electron transfer kinetics between electrode and protein. The redox center is usually buried inside the protein, and therefore the tunneling probability for the electron from the electrode surface to the redox center is low. Moreover, as proteins are macromolecules, their diffusion into the double layer is slower than for small molecule mediators. And finally, if the electrode surface is covered with the enzyme, for example, as in biosensors, the number of redox centers per electrode area is rather low and the current densities (i.e., productivities) of the reactions are rather too low for synthetic purposes. A variety of different mediator types, such as metal complexes or aromatic compounds, have therefore been investigated and engineered to improve the reaction. In comparison with established biotransformation processes (e.g., whole cell reductions or enzyme coupled cofactor regeneration) (see Figure 39.2), the productivity of the majority of electroenzymatic processes is rather low. In addition, the final product concentration is often too low for a practical synthesis [22]. The heterogeneous electrochemical reaction and the enzymatic process have to be carefully analyzed to identify the rate-limiting step. The overall system needs to be optimized by means of reaction engineering. Acceleration of the electrochemical reaction is in principle relatively easy to achieve and several approaches have been demonstrated already. First, a 3D electrode design (e.g., a packed bed of graphite particles) allows establishing a large electrode surface/ volume ratio. In 3D cells, productivities of up to 1 kg L−1 d−1 of reduced cofactor [95] or 100 g L−1 d−1 of chiral product (Table 39.6, entry 10) have been achieved. When working with such high productivities, the medium composition has to be carefully considered to establish a sufficiently high conductivity, for example, by adding salts or ionic liquids as supporting electrolyte [67]. In these cases, the biocompatibility of the electrolyte has to be checked. But other parameters like the mediator concentration have to be carefully adjusted with respect to the enzyme reaction performance. Low final product concentrations can be circumvented by using an aqueous-organic two-phase system [46,52] to overcome the solubility limitations for hydrophobic substances in the water phase. Again, ionic liquids can be used as water-miscible additives to increase substrate solubility synergistically in addition to their increase of conductivity. Electrochemistry provides an elegant method for controlled supply of (co)substrates. For biotransformations with peroxidases, which require H2O2 as cosubstrate, several successful examples can be found with in situ electrogeneration of the hydrogen peroxide. The control of the potential was used as a convenient tool for adjusting the production rate. Paired electrolysis, which combines

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Electroenzymatic Synthesis

1539

electrochemical cofactor regeneration with reactant supply, is in principle a very smart concept and the proof of concept was demonstrated with a few systems. Additional opportunities to establish these systems can be foreseen. Improving the reaction productivity remains the main challenge in order to benefit from the advantages of electrochemistry in biocatalysis. Nearly all the processes described were established without focusing on the electrochemical cell design. Thus, reactor concepts with enhanced mass transfer like 3D cells have by no means been exhausted for electroenzymatic processes and offer a chance for improved performance. It remains an open question whether electrochemistry can compete with other methods of supplying redox equivalents for enzymatic biotransformations. Some advantages have not been fully exploited due to a lack of data (e.g., full life cycle assessment and comparison of waste) between a fully enzymatic process and an electrochemical process for the same product. Some new challenges, like the detailed understanding of the inactivation mechanism between the otherwise successful rhodium mediator and proteins, have only recently been identified. Moreover, promising concepts like the use of ionic liquids and optimized electrochenmical cell design have not yet been combined. Therefore, this area of research—the combination of enzyme catalysis and electrochemistry—remains an open field and, as Steckhan described it, “only the tip of the iceberg has surfaced” [19].

REFERENCES 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28.

Drauz, K.; Waldmann, H. Enzyme Catalysis in Organic Synthesis; VCH: Weinheim, Germany, 1995. Faber, K. Biotransformation in Organic Chemistry; 5th edn.; Springer-Verlag: Berlin, Germany, 2004. Bornscheuer, U. T.; Buchholz, K. Eng. Life Sci. 2005, 5, 309–323. Liese, A.; Seelbach, K.; Wandrey, C. Industrial Biotransformations; 2nd edn.; Wiley-VCH: Weinheim, Germany, 2006. Schmid, A.; Hollmann, F.; Park, J. B.; Buhler, B. Curr. Opin. Biotechnol. 2002, 13, 359–366. Reetz, M. T. Angew. Chem. Int. Ed. 2011, 50, 138–174. Turner, N. J. Trends Biotechnol. 2003, 21, 474–478. Lütz, S.; Giver, L.; Lalonde, J. Biotechnol. Bioeng. 2008, 101, 647–653. Webb, E. C. Enzyme Nomenclature; International Union of Biochemistry and Molecular Biology: Academic Press Inc., San Diego, CA, 1992. Goldberg, K.; Schroer, K.; Lütz, S.; Liese, A. Appl. Microbiol. Biotechnol. 2007, 76, 249–255. Goldberg, K.; Schroer, K.; Lütz, S.; Liese, A. Appl. Microbiol. Biotechnol. 2007, 76, 237–248. Wandrey, C. Chem. Rec. 2004, 4, 254–265. Wichmann, R.; Vasic-Racki, D. In: Technology Transfer in Biotechnology: From Lab to Industry to Production; Kragl, U.; Scheper, T. (eds.) Vol. 92; Springer: Berlin, Germany, 2005, pp. 225–260. Lütz, S. In: de Vries, J. G., Elsevier, C. J. (eds.) The Handbook of Homogeneous Hydrogenation; Vol. III; Wiley-VCH Verlag GmbH: Weinheim, Germany, 2006, pp. 1471–1482. Hollmann, F.; Schmid, A. Biocatal. Biotrans. 2004, 22, 63–88. Hummel, W.; Kula, M. R. Eur. J. Biochem. 1989, 184, 1–13. Findrik, Z.; Vasic-Racki, D.; Lütz, S.; Daussmann, T.; Wandrey, C. Biotechnol. Lett. 2005, 27, 1087–1095. Hildebrand, F.; Lütz, S. Tetrahedron Asymm. 2006, 17, 3219–3225. Steckhan, E. Top. Curr. Chem. 1994, 170, 83–111. Steckhan, E.; Arns, T.; Heineman, W. R.; Hilt, G.; Hoormann, D.; Jörissen, J.; Kroner, L.; Lewall, B.; Pütter, H. Chemosphere 2001, 43, 63–73. Wienkamp, R.; Steckhan, E. Angew. Chem. Int. Ed. Engl. 1982, 21, 782–783. Ruinatscha, R.; Höllrigl, V.; Otto, K.; Schmid, A. Adv. Synth. Catal. 2006, 348, 2015–2026. Kohlmann, C.; Märkle, W.; Lütz, S. J. Mol. Catal. B: Enzym. 2008, 51, 57–72. Steckhan, E. Angew. Chem. Int. Ed. Engl. 1986, 25, 683–701. Bonnefoy, J.; Moiroux, J.; Laval, J. M.; Bourdillon, C. J. Chem. Soc. Faraday Trans. 1 1988, 84, 941–950. Hilt, G.; Lewall, B.; Montero, G.; Utley, J. H. P.; Steckhan, E. Liebigs Ann./Recl. 1997, 11, 2289–2296. Hilt, G.; Steckhan, E. J. Chem. Soc.-Chem. Commun. 1993, 1706–1707. Komoschinski, J.; Steckhan, E. Tetrahedron Lett. 1988, 29, 3299–3300.

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29. Degenring, D.; Schröder, I.; Wandrey, C.; Liese, A.; Greiner, L. Org. Process Res. Dev. 2004, 8, 213–218. 30. Schröder, I.; Steckhan, E.; Liese, A. J. Electroanal. Chem. 2003, 541, 109–115. 31. Manjon, A.; Obon, J. M.; Casanova, P.; Fernandez, V. M.; Ilborra, J. L. Biotechnol. Lett. 2002, 24, 1227–1232. 32. Schulz, M.; Leichmann, H.; Gunther, H.; Simon, H. Appl. Microbiol. Biotechnol. 1995, 42, 916–922. 33. Itoh, S.; Fukushima, H.; Komatsu, M.; Ohshiro, Y. Chem. Lett. 1992, 1583–1586. 34. Kashiwagi, Y.; Osa, T. Chem. Lett. 1993, 677–680. 35. Wulf, H.; Perzborn, M.; Sievers, G.; Scholz, F.; Bornscheuer, U. T. J. Mol. Catal. B: Enzym. 2012, 74, 144–150. 36. Burnett, J. N.; Underwood, A. L. Biochemistry 1965, 4, 2060–2064. 37. Powning, R. F.; Kratzing, C. C. Arch. Biochem. Biophys. 1957, 66, 249–251. 38. Kono, T. Bull. Agric. Chem. Soc. Jpn. 1957, 21, 115–120. 39. Kono, T.; Nakamura, S. Bull. Agric. Chem. Soc. Jpn. 1958, 22, 399–403. 40. Baik, S. H.; Kang, C.; Jeon, I. C.; Yun, S. E. Biotechnol. Tech. 1999, 13, 1–5. 41. Siu, E.; Won, K.; Park, C. B. Biotechnol. Prog. 2007, 23, 293–296. 42. Steckhan, E. In: Lund, H., Hammerich, O. (eds.) Organic Electrochemistry; 4th edn.; Marcel Dekker: New York, 2000. 43. Ruppert, R.; Herrmann, S.; Steckhan, E. Tetrahedron Lett. 1987, 28, 6583–6586. 44. Delecouls-Servat, K.; Basseguy, R.; Bergel, A. Chem. Eng. Sci. 2002, 57, 4633–4642. 45. Delecouls-Servat, K.; Basseguy, R.; Bergel, A. 2002, 55, 93–95. 46. Hildebrand, F.; Lütz, S. Tetrahedron Asymm. 2007, 18, 1187–1193. 47. Hollmann, F.; Witholt, B.; Schmid, A. J. Mol. Catal. B: Enzym. 2002, 19, 167–176. 48. Grau, M. M.; Poizat, M.; Arends, I. W. C. E.; Hollmann, F. Appl. Organomet. Chem. 2010, 24, 380–385. 49. Hildebrand, F.; Kohlmann, C.; Franz, A.; Lütz, S. Adv. Synth. Catal. 2008, 350, 909–918. 50. Lutz, J.; Hollmann, F.; Ho, T. V.; Schnyder, A.; Fish, R. H.; Schmid, A. J. Organomet. Chem. 2004, 689, 4783–4790. 51. Hildebrand, F.; Lütz, S. Chem. Eur. J. 2009, 15, 4998–5001. 52. Höllrigl, V.; Otto, K.; Schmid, A. Adv. Synth. Catal. 2007, 349, 1337–1340. 53. Kashiwagi, Y.; Yanagisawa, Y.; Shibayama, N.; Nakahara, K.; Kurashima, F.; Anzai, J.; Osa, T. Electrochim. Acta 1997, 42, 2267–2270. 54. Kang, Y. W.; Kang, C.; Hong, J. S.; Yun, S. E. Biotechnol. Lett. 2001, 23, 599–604. 55. Kim, M. H.; Yun, S. E. Biotechnol. Lett. 2004, 26, 21–26. 56. DiCosimo, R.; Wong, C. H.; Daniels, L.; Whitesides, G. M. J. Org. Chem. 1981, 46, 4622–4623. 57. Yuan, R.; Watanabe, S.; Kuwabata, S.; Yoneyama, H. J. Org. Chem. 1997, 62, 2494–2499. 58. Cantet, J.; Bergel, A.; Comtat, M. Enzyme Microb. Technol. 1996, 18, 72–79. 59. Massey, V. J. Biol. Chem. 1994, 269, 22459–22462. 60. Trost, E.-M.; Fischer, L. J. Mol. Catal. B: Enzym. 2002, 19–20, 189–195. 61. Butó, S.; Pollegioni, L.; D´Angiuro, L.; Pilone, M. S. Biotechnol. Bioeng. 1994, 44, 1288–1294. 62. Armstrong, F. A.; Hill, H. A. O.; Walton, N. J. Acc. Chem. Res. 1988, 21, 407–413. 63. Hill, H. A. O.; Oliver, B. N.; Page, D. J.; Hopper, D. J. J. Chem. Soc.-Chem. Commun. 1985, 1469–1471. 64. Frede, M.; Steckhan, E. Tetrahedron Lett. 1991, 32, 5063–5066. 65. Brielbeck, B.; Frede, M.; Steckhan, E. Biocatalysis 1994, 10, 49–64. 66. Petersen, A.; Steckhan, E. Bioorg. Med. Chem. 1999, 7, 2203–2208. 67. Kohlmann, C.; Greiner, L.; Leitner, W.; Wandrey, C.; Lütz, S. Chem. Eur. J. 2009, 15, 11692–11700. 68. van Rantwijk, F.; Sheldon, R. A. Chem. Rev. 2007, 107, 2757–2785. 69. Kragl, U.; Eckstein, M.; Kaftzik, N. Curr. Opin. Biotechnol. 2002, 13, 565–571. 70 Jeon; Jin, S.; Shin, I. H.; Sang, B. I.; Park, D. H. J. Microbiol. Biotechnol. 2005, 15, 281–286. 71. Hollmann, F.; Hofstetter, K.; Schmid, A. Trends Biotechnol. 2006, 24, 163–171. 72. Bernhardt, R. J. Biotechnol. 2006, 124, 128–145. 73. Hannemann, F.; Bichet, A.; Ewen, K. M.; Bernhardt, R. Biochim. Biophys. Acta, Gen. Subj. 2007, 1770, 330–344. 74. Faulkner, K. M.; Shet, M. S.; Fisher, C. W.; Estabrook, R. W. Proc. Natl. Acad. Sci. USA 1995, 92, 7705–7709. 75. Udit, A. K.; Arnold, F. H.; Gray, H. B. J. Inorg. Biochem. 2004, 98, 1547–1550. 76. Hollmann, F.; Schmid, A.; Steckhan, E. Angew. Chem. Int. Ed. 2001, 40, 169–171. 77. Hollmann, F.; Hofstetter, K.; Habicher, T.; Hauer, B.; Schmid, A. J. Amer. Chem. Soc. 2005, 127, 6540–6541. 78. Kuwabata, S.; Morishita, N.; Yoneyama, H. Chem. Lett. 1990, 19, 1151–1154.

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79. Kawabata, S.; Iwata, N.; Yoneyama, H. Chem. Lett. 2000, 29, 110–111. 80. Seelbach, K.; vanDeurzen, M. P. J.; vanRantwijk, F.; Sheldon, R. A.; Kragl, U. Biotechnol. Bioeng. 1997, 55, 283–288. 81. van de Velde, F.; van Rantwijk, F.; Sheldon, R. A. J. Mol. Catal. B: Enzym. 1999, 6, 453–461. 82. Pletcher, D. Acta Chem. Scand. 1999, 53, 745–750. 83. Bartlett, P. N.; Pletcher, D.; Zeng, J. J. Electrochem. Soc. 1999, 146, 1088–1092. 84. Chen, J. K.; Nobe, K. J. Electrochem. Soc. 1993, 140, 299–303. 85. Lee, K.; Moon, S. H. J. Biotechnol. 2003, 102, 261–268. 86. Dembitsky, V. M. Tetrahedron 2003, 59, 4701–4720. 87. Laane, C.; Weyland, A.; Franssen, M. Enzyme Microb. Technol. 1986, 8, 345–348. 88. Lütz, S.; Steckhan, E.; Liese, A. Electrochem. Commun. 2004, 6, 583–587. 89. Kohlmann, C.; Lütz, S. Eng. Life Sci. 2006, 6, 170–174. 90. Lütz, S.; Vuorilehto, K.; Liese, A. Biotechnol. Bioeng. 2007, 98, 525–534. 91. Kuwabata, S.; Watanabe, S.; Inoue, H.; Yoneyama, H. Denki Kagaku 1996, 64, 1080–1083. 92. Biade, A. E.; Bourdillon, C.; Laval, J. M.; Mairesse, G.; Moiroux, J. J. Amer. Chem. Soc. 1992, 114, 893–897. 93. Anne, A.; Bourdillon, C.; Daninos, S.; Moiroux, J. Biotechnol. Bioeng. 1999, 64, 101–107. 94. Märkle, W.; Lütz, S. Electrochim. Acta 2008, 53, 3175–3180. 95. Vuorilehto, K.; Lütz, S.; Wandrey, C. Bioelectrochemistry 2004, 65, 1–7.

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40

Electrochemical Modeling of Biological Processes Richard D. Webster

CONTENTS I. II.

Introduction ........................................................................................................................ 1543 Electrochemistry Coupled to Liquid Chromatography–Mass Spectrometry ..................... 1544 A. Experimental Considerations ...................................................................................... 1544 B. Comparison of Electrochemical Oxidation with Cytochrome P450 Oxidation ......... 1545 C. Examples of Metabolic Reaction Products Detected by EC–LC–MS ....................... 1546 1. Clozapine ............................................................................................................. 1546 2. Paracetamol.......................................................................................................... 1546 3. Amodiaquine ....................................................................................................... 1548 4. Diclofenac ............................................................................................................ 1548 5. Lidocaine ............................................................................................................. 1549 6. Boscalid ............................................................................................................... 1549 7. Toremifene ........................................................................................................... 1550 8. EC–LC Coupled to ICP–MS................................................................................ 1551 D. Polycyclic Aromatic Hydrocarbons ............................................................................ 1551 III. Cyclic Voltammetry Studies of Low-Molecular-Weight Molecules Involved in Biological Reactions ........................................................................................................... 1552 A. Vitamin A ................................................................................................................... 1552 B. Vitamin B2 .................................................................................................................. 1553 C. Vitamin E .................................................................................................................... 1555 D. Vitamin K1 .................................................................................................................. 1558 E. Dopamine.................................................................................................................... 1560 F. Voltammetry of Low-Molecular-Weight Molecules in Model Biological Membranes .... 1561 IV. Concluding Remarks .......................................................................................................... 1562 References .................................................................................................................................... 1563

I. INTRODUCTION There are many biological processes that involve reduction or oxidation (redox) steps, including enzymatic reactions, photosynthesis reactions, free radical oxidative reactions, and metabolite processes. While it is possible in some circumstances to use electrochemical methods to monitor individual chemicals that exist within biological systems (such as dopamine (DA), which is present as a neurotransmitter in a wide variety of animals), it is very difficult to study complex redox reaction mechanisms that occur in vivo, due to the many interfering species that are also present. However, electrochemical methods are extremely useful in studying the exact pathways that the redox processes occur by, usually by isolating the individual chemicals and performing detailed experiments under carefully controlled laboratory conditions. Knowledge of the exact reaction mechanism, including the number of electrons involved and the order of electron transfer, can be critical in identifying

1543

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Organic Electrochemistry

intermediates that are also likely to exist in biological systems. In this chapter, the discussion focuses on relatively low-molecular-weight ( benzene (Bz). As the nucleophilicity of the solvents commonly used in electropolymerization rises in the order HF < CF3COOH < SO2 < CH3NO2 < CH2Cl2 < propylene carbonate (PC) < CH3CN < H2O, only Py can be polymerized in water [24]. Even in the case of Th, small amounts of water are sufficient for nucleophilic addition to block further growth of the oligomer chain [25]. By contrast, Bz can be only partially polymerized in CH3CN or PC, as the newly formed oligomers react with the solvent more readily than with the cations of the oligomeric intermediates. In dry SO2, conductive polyparaphenylene films are formed in high yield [26]. Recently, ionic liquids (ILs) that possess a very low nucleophilicity (see Chapter 8) and a wide potential window (4–6 V) have facilitated the synthesis of CPs [27]. Thus, all the classic CPs have been prepared in ILs, for example, PP [28] and PTh [29]. A disadvantage is their high viscosity, which diminishes the diffusion of electroactive species to the electrode and may be unfavorable for applications. Another solvent/electrolyte system that has been used for the electrosynthesis of CPs is boron trifluoride diethyl ether (BFEE) [30]. Its main characteristic is the considerable lowering of the oxidation potential that is necessary for a successful electropolymerization. Numerous CPs have been prepared in the presence of BFEE. In particular, this strategy has been applied to many of those monomers that are difficult to oxidize, for example, benzene [31], furan [32], anthracene [33], or 5-bromo-indole [34].

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Organic Electrochemistry

The new trend to “green” chemistry has stimulated the electropolymerization in microemulsion (see Chapter 8) media [35]. Microemulsions are systems that at least consist of three components: water, a hydrophobic organic material (starting monomer), and an amphiphilic surfactant. By mixing these components, the solubility of water-insoluble monomers has been drastically improved. Thus, the electropolymerization of EDOT in water has been successfully carried out in the presence of the nonionic polyoxyethylene-10-lauryl ether [36]. In 1981, Diaz et al. [10] suggested, by analogy to the long-known coupling reactions of radical cations in aromatic compounds, that in the polymerization of Py, the monomers dimerize at the α-position after oxidation at the electrode and that protons are eliminated from the doubly charged dihydrodimer, forming the neutral species. As the neutral dimer, on account of its greater conjugation, is more easily oxidized than the monomer under the given experimental conditions, it is immediately oxidized to its cation. Chain growth should be accompanied by the addition of new cations of the monomeric Py to the already charged oligomers. This, in turn, is followed by another proton elimination and the oxidation of the propagated oligomeric unit to a cation (Scheme 41.1). This classic chain propagation mechanism is still widely accepted in the literature [5b,37]. However, recent studies using the so-called oligomer approach [38] have shown that this mechanistic view is far too simple. Obviously, after the formation of a dimer, a sequence of subsequent dimerization steps leads to the formation of soluble oligomers with chain lengths normally ranging between four and eight units [39]. All these reactions preferably occur in solution without or with only small precipitation on the electrode. Subsequently, deposition and growth processes set in, triggered by nucleation reactions. They are strongly influenced by the concentration of the starting monomer, the temperature [40], and relevant electrochemical parameters such as the formation potential [41]. Especially, at high concentrations of monomer, the deposition may begin at a shorter chain length. Recent studies have also shown that coupling reactions between charged oligomers may occur also with dicationic species, which enlarges the reactivity of oligomers [42]. Quantitative investigations of the kinetics of these α-coupling steps suffered because rate constants were beyond the timescale of voltammetric experiments until ultramicroelectrodes and improved electrochemical equipment made possible a new transient method called fast-scan voltammetry [43]. With this technique, cyclic voltammetric experiments up to scan rates of 1 MV s−1 are possible, and species with lifetimes in the nanosecond scale can be observed. Using this technique, Hapiot et  al. [44] were the first to obtain data on the lifetimes of the electrogenerated Py radical cation and substituted derivatives. The resulting rate constants for the dimerization of such monomers lie in the order of 109 M−1s−1. The same authors were also able to show that the radical cation of the Py tetramer, by contrast, is rather stable [44]. This tendency of decreasing reactivity as a function of chain length is a general property of all chainlike conjugated oligomers [45]. Thus, in the series of unsubstituted thiophenes, the radical cations of monomeric thiophene dimerize with a rate constant greater than 109 M−1s−1, while the lifetimes of oligomer cation radicals increase with chain length [46]. For example, the rate constant for the dimerization of the unsubstituted thiophene tetramer is about 105 M−1s−1 [46a,47], and that of the pentamer is below 104 M−1s−1 [48]. +

–2e

2

+

+

X

X

X

X

X

X

X

n X = NH, S

Classical formation mechanism of CPs.

–2H+

X X

+

X

X

+

X

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H

+

+

X

SCHEME 41.1

H

+

X

X n

+ 2H+

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Electrochemistry of Conducting Polymers

X

2

X X

+

–2e

2

–2e E 02

E 01

2

X

+

+

X

kf

≈ 109

X

kf ≈ 108

X

X X

+

H X

+

–2H+

X

H

X X

H

+

+

+

–2H+

X

X

H

+

X

´

2

X X

X n=2

E 03

2

+

X X

X n=2

X

+ H X X + n=2 H X = NH, S

X

–2H+ kf ≤ 1

X n=2

H

X X n=2

X H +

3,3´

X

X

X n=2

kf ≤ 104

X

e.g.,

X

X X + n=2 H

X

2,2

–2e

X

kf < 10

2

X X n=2

+ H X n=2

X X

X n=6

Deposition

E 01 > E 02 > E 03

SCHEME 41.2 Initial steps of electropolymerization involving σ-bonded intermediates.

From all these measurements, it is clear that the rate constants of the dimerization of chainlike conjugated oligomers, and of their coupling steps with the original monomer, decrease with increasing chain length. In addition, the rates of all these second-order reactions are a function of the concentrations of the reacting species. Therefore, as can be shown by simulations, a monomeric radical cation as starting species is far more likely to couple with monomeric radical cations than with radical cations of higher oligomers (Scheme 41.2). Studies by Heinze et al. on donor-substituted thiophenes [49] such as methylthio (= methylsulfanyl) or methoxy-substituted derivatives provide further clear evidence for this reaction pathway. They found, for instance, that 3-methylthiothiophene 1 undergoes a fast coupling reaction. However, deposition processes or insoluble film formation could not be detected in any experiments with this compound under any conditions, even at high concentrations. Similarly, the 3,3′-disubstituted bithiophene 2a does not polymerize, but the anodic oxidation of 4,4′-dimethylthiobithiophene 2 produces an excellent yield of CP. A careful analysis based on these experimental results excluded a chain propagation process [49a]. On account of the third position of the methylthio substituent in the thiophene ring, three isomeric dimers may be formed. The main reaction path can be deduced from the mesomeric forms of the radical cation 1+ •. The two most important mesomeric structures are those with the unpaired electron in an α-position. The mesomeric structure with the positive charge next to the methylthio sulfur is preferred, because of the stabilizing +M-effect of the methylthio substituent. Therefore, the 2,2′-connected dimer 2a is the essential product of the radical coupling process. Dimers 2b and 2 are minor side products (Scheme 41.3). The main product initially formed, dimer 2a, will undergo further slow coupling steps. At the given potential, dimer 2a will immediately be oxidized to the corresponding mono-radical cation. Again, the mono-radical cation can be expressed in different mesomeric structures. The most reasonable notation has the positive charge next to the methylthio sulfur and the unpaired electron at the blocked inner α-position:

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1576

Organic Electrochemistry Fast

R +

S

R

*2

S

–2H+

2a

S R

R

R

–e– +e–

S

S

–2H+

2b

S R

1 R

R *2

S

+

S

–2H+

2

S R

Slow R = –SCH3 –OCH3

SCHEME 41.3

First coupling steps after the oxidation of 3-methylthiothiophene or 3-methoxythiophene. +

SCH3 +

SCH3 S

S S

S

CH3S

CH3S

Most important mesomeric structures of the radical cation of 2a

Therefore, the reactivity of this species, similar to that of the 3,3′-dimethoxy-2,2′-bithiophene radical cation, is low, probably resulting in a rate constant for the dimerization of 100 mV s−1), broad waves are observed during the cathodic reverse scan at potentials around 0 V. This is typical for the discharging of protons formed during the process. The coulometric analysis of the voltammograms shows that one charge is lost per molecule (by proton cleavage) in the condensation reaction. The average functionality of a monomer unit, f (= number of reactive sites of a monomer that participate in polymerization), has been calculated from coulometric data [88]. The resulting values of f < 2, together with all other observations, give clear evidence that the short-chain H–T8–H dimerizes quantitatively in

(a)

(b)

(c)

0.0

0.5 1.0 E (vs. Ag/AgCl) (V)

FIgURE 41.5 Potentiodynamically generated solid-state coupling of octathiophene, v = 10 mV s−1, T = −5°C: (a) cyclic voltammogram of octathiophene, Eλ = 1.07 V; (b) cyclic voltammetry during coupling, Eλ = 1.25 V; (c) cyclic voltammogram of sedecimthiophene, Eλ = 1.25 V. The resulting sedecimthiophene forms a stable σ-interchain product after oxidation. (Reprinted from Electrochim. Acta, 41, Meerholz, K. and Heinze, J., 1839, Copyright (1996), with permission from Elsevier.)

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Electrochemistry of Conducting Polymers

1583

this solid-state reaction, forming an isomer of sedecimthiophene. Analogous reactions have been observed for other short-chain oligomers, such as sexithiophene or sexiphenylene, leading to dodecathiophene and -phenylene. Of particular note is the strong hysteresis between the charging and discharging steps during the solid-state polymerization of all these oligomers (Figure 41.5c). This phenomenon is characteristic of the formation and the decay of intermediate σ-dimers, the stability of which increases as a function of chain length. All these materials can be polymerized further at higher formation potentials. The number of coupling steps depends strongly on the applied potential. At the end of such processes, the voltammograms have the typical shape of those of CPs in general, that is, they exhibit the characteristic current plateau and high current waves at the onset of charging and the end of discharging. Under these conditions, the average functionality may be larger than the limiting value for infinite chains (f = 2). From this, however, it must be concluded that chain-lengthening steps as well as coupling reactions between the chains take place, leading to a network with an intact π-system. All these experiments give no indication of a polymer of infinite length. Based on IR data, Furukawa et al. concluded from studies of thiophene oxidation that the degree of polymerization is rather low, ranging from 10 to 35 [89]. For the conjugation length of PTh, a number of approximately 11 has been estimated [90]. In summary, all the recent findings published in the literature exclude a simple chain propagation mechanism for the formation of CPs. Rather, electropolymerization involves three different stages: 1. Oxidation of the monomer at the electrode, formation of soluble oligomers in the diffusion layer, preferably successive dimerization step, autocatalytic reactions 2. Deposition of oligomers, involving nucleation and growth processes 3. Solid-state polymerization, producing longer chains and cross-linked materials

C.

EXPERIMENTAL TECHNIQUES AND CONDITIONS

A well-known phenomenon of CPs is their challenging diversity. This is induced by even small variations of experimental parameters such as formation potential, concentration of the starting monomer, or temperature creating significant changes of physical and chemical properties. The essential reason for this complex behavior is based on the fact that each coupling step requires electrochemical activation. Therefore, already the specific (respective) electrochemical preparation technique influences the results of electropolymerization experiments the most important of which are potentiodynamic, potentiostatic, and galvanostatic polymerization [74]. Potentiodynamic polymerization (cyclic voltammetry) is characterized by a cyclic regular change of the electrode potential during the deposition of the CP onto the electrode. The growing polymer film—following the potential changes—continuously changes between its neutral (insulating) and its doped (conducting) states, which are accompanied by a continuous exchange of electrolyte and solvent through the freshly deposited polymer. This automatically provokes changes in the polymer matrix and favors the formation of disordered chains. Typically, CPs generated by potentiodynamic polymerization are obtained in their neutral state at the end of the preparation. Conversely, potentiostatic or galvanostatic techniques lead to materials in their doped state. A significant advantage of these static methods of electropolymerization is that well-ordered structures are prepared, which do not experience any changes after deposition and succeeding solid-state reactions. A characteristic example for this phenomenon is PPy that has been galvano- or potentiostatically polymerized at low formation potentials in PC/0.1 M LiClO4. Surprisingly, during the very first voltammetric discharge cycle, the reduction is significantly shifted to negative potentials, and a strong mass increase proves an efficient cation insertion despite the fact that anions are relatively small [91–93]. At low temperatures, this cation insertion completely dominates and is accompanied by a strong solvent insertion. In subsequent cycles, the cation effect and the correlated reduction shift significantly diminish (Figure 41.6). According to the mechanistic pattern presented by Heinze [94], the potentio- or galvanostatic formation of CPs ends at the level of well-ordered σ-dimers, the

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Organic Electrochemistry 150

Current μA

100 50 0 –50 –100 –150 –1.5

1.0

2.0 – 4.0

–1.0

–0.5 0 E (V) vs. Ag/AgCl

0.5

1.0

–1.0

–0.5 0 E (V) vs. Ag/AgCl

0.5

1.0

(a)

0.4

Frequency ( kHZ)

0.2

0.0

–0.2

–0.4

–0.6 –1.5 (b)

FIgURE 41.6 Multisweep cyclic voltammetry and EQCM measurements of a PPy film after potentiostatic polymerization (Ep = 0.8 V) of Py in PC (c = 10 mM, c(LiClO4) = 0.1 M); (a) CV v = 20 mV s−1, first up to fourth cycle after polymerization, (b) frequency—potential plot (EQCM) (From Bilger, R., Dissertation, Freiburg, Germany, 1994.)

first discharge of which occurs via an efficient cation insertion. The reason for this behavior results from the fact that the positive charges in σ-dimers are not mobile. Besides the electrochemical method applied during polymerization, experimental parameters such as potential, temperature, and solvent are also important for the quality and the properties of CPs. Especially, the formation potential determines the chain length and structure of CPs. Very high oxidation potentials or currents in galvanostatic experiments, which imply the generation of highly charged and reactive intermediates, lead to defects and the formation of cross-linked materials [86]. At low potentials or currents, the oligomeric intermediates are weakly charged, and consequently the electropolymerization may end at an oligomeric level with chain lengths between 10 and 20 units [41,95]. Again, the electropolymerization of Py documents the influence of different formation potentials. Figure 41.7 shows multisweep voltammograms of PPy films that have been generated at positive switching potentials between 0.83 and 1.13 V [52,96]. As can be seen, in addition to the normal wave of PPy at Ep ~ 0.01 V vs. Ag/AgCl, a new sharp oxidation wave emerges at potentials about

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Electrochemistry of Conducting Polymers

4 μA

–1.2 –0.8 –0.4 0.0 0.4 0.8 1.2

0.1 μA 0.1 μA

–1.2 –0.8 –0.4 0.0

0.4

0.8

1.2

–1.2 –0.8 –0.4 0.0

0.4

0.8

1.2

FIgURE 41.7 Potential dependence of potentiodynamic growth of PPy films. Nondegassed acetonitrile solution, 0.1 M Py, 0.1 M TBAPF6, 1 wt% water, T = −20°C, v = 100 mV s−1. (Left) −1.12 to 1.13 V, 15 scans; (center) −1.12 to 0.88 V, 28 scans; (right) −1.12 to 0.83 V, 30 scans and scan range changed to −1.12 to 0.78 V for another 28 scans. (Reprinted from Electrochim. Acta, 44, Zhou, M. and Heinze, J.,1733, Copyright (1999), with permission from Elsevier.)

−0.23 V when the voltammetric switching potential is 0.83 V. The same observations result with PPy synthesized at extremely low currents with the galvanostatic technique [95]. These data in combination with spectroscopic findings provide evidence that at least two types of PPy exist, the so-called PPy-I with chain lengths up to 64 units and PPy-II with chain lengths between 12 and 16 units [41]. A cross-linked material referred to as Py-III is generated at potentials higher than 1 V. The second most important parameter during electropolymerization is the temperature. Its influence on the kinetics and the resulting properties of CPs is considerable. As already described, the polymerization is a multistep process in which each sub-step possesses its own activation energy. Therefore, at low temperatures, the rates of coupling steps and proton eliminations decrease. On the other side, the thermodynamic equilibria between charged intermediates and corresponding dimers shift in favor of the coupling products due to entropic reasons and decreasing coulombic repulsions between charged moieties [57a]. Thus, σ-intermediates could be detected [97], and systems with shorter chain lengths are favored [52,96]. Obviously, low temperatures stabilize well-ordered structures and enhance conductivities of CPs [40,98,99]. For example, conductivities of PPy films generated at 234 K reaches conductivity values higher than 1000 S cm−1 [100]. Moreover, despite the fact that the reaction rates slow down, the yield of deposited materials may increase because the solubility of oligomers decrease [83]. Other factors that influence the polymerization process and the final properties of the generated CPs are solvents, additives (e.g., water, bases, acids), and electrolytes. In order to get an efficient polymer yield, solvents should possess a high polarity, which minimizes coulombic repulsions during cationic coupling steps [61]. On the other hand, their nucleophilicity should be low. In the case of conventional solvents, it rises in the following order: CH3NO2 < CH2Cl2 < PC < CH3CN < H2O. Therefore, water can be used only for the generation of polymers for which the oxidation potentials of their monomers are low, for example, Py or EDOT. Nevertheless, even for these systems, the electropolymerization process produces many defects due to nucleophilic attacks of water to the cationic intermediates [101]. The most popular organic solvent is acetonitrile. However, it is not the best one for the preparation of CPs. Due to its low basicity—the pKBH+ value of protonated acetonitrile is about −10 [102]—superdry acetonitrile does not polymerize Py. The reaction stops at the level of weakly acidic σ-intermediates of bi- or more likely tetrapyrrole. A stronger base than acetonitrile must be used to initiate the elimination of protons. Water fulfills this condition. Py can be electropolymerized in acetonitrile in the presence of 1% water [9b,103]. A similar effect results from the application of a sterically hindered base such as 2,6-di-tert-butylpyridine [82]. Solvents with a weak nucleophilicity, which can be also applied for monomers with high oxidation potentials such as benzene, are CF3COOH, HF, BF3-etherate, SO2, and ILs.

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STRUCTURE AND MORPHOLOgy OF CONDUCTINg POLyMERS

Despite the fact that CPs are semiconductors, their similarities to conventional inorganic semiconductors are superficial. The most prominent difference concerns the structure. While inorganic semiconducting materials normally possess a periodical lattice structure with a high degree of 3-D order, electrochemically fabricated CPs are typically strongly disordered involving chainlike structures on the molecular level. Consequently, the electronic properties of inorganic systems are described on the basis of band models, whereas the picture of CPs is currently changing by a stronger focus on molecular characteristics beyond the original classification as organic solid-state bodies with “infinite” chain lengths. Besides synthesis, current basic research on CPs is concentrated on structural analysis. Structural parameters—for example, not only regularity and homogeneity of chain structures but also chain length—play an important role in our understanding of the properties of such materials and are important for applications. Research on electropolymerized materials has concentrated on PPy and PTh derivatives in particular and PANI as well [8]. In the last few decades, the oligomer approach has been successfully applied to the study of physical and chemical properties of CPs [38]. When monitoring the physical properties as a function of chain length, extrapolation leads to the properties of a defect-free polymer. Spectroscopic methods have proved particularly suitable for characterizing structural properties. These comprise surface techniques such as XPS, AES, or ATR, on the one hand, and the usual methods of structural analysis, such as NMR, ESR, and x-ray diffraction techniques, on the other hand. PPy was the first CP to be structurally analyzed. The discovery that α,α′-disubstituted Py did not polymerize led to the conclusion that the Py units in PPy are α-linked [104]. Magic angel spinning 13C NMR data support the view that the Py units bond chiefly in the α,α′-position, although α,β bondings are also found [105]. On the other hand, XPS measurements of PPy reveal that a full onethird of the Py rings in a chain are irregularly bonded [106]. Very recent electrochemical studies have shown that the structure of PPy strongly depends on the preparation conditions during electropolymerization [52,96,103]. In highly acidic solutions, saturated pyrrolidine units are generated, which lead to the formation of a passivating film. At very positive formation potentials greater than 1.2 V vs. Ag/AgCl, a cross-linked material is generated, while at low oxidation potentials, relatively short linear chains are produced. Similarly, a chain structure with predominantly α,α′-coupling between the monomer units is postulated for PTh on the basis of spectroscopic findings and mechanistic studies. IR measurements of uncharged chemically produced examples in particular clearly reveal that α,α′-bonding predominates in the polymers produced from α,α′-dibromothiophene, 2,2′-bithiophene, or 3-methylthiophene [4,107]. The IR data correlate very well with 13C NMR measurements, which similarly confirm the dominance of α,α′-bonding in P3-MeTh [108]. In agreement with these findings, the electrochemical oxidation of 3,3′,5,5′-tetramethyl-2,2′-bithiophene unambiguously demonstrates that the nonblocked outer β-position is nonreactive [57c]. Chain length is another factor closely related to the structural characterization of CPs. The importance of this parameter lies in the considerable influence of the electric as well as the electrochemical properties of CPs. However, the molecular weight techniques normally used in polymer chemistry cannot be employed on account of extreme insolubility of the materials. A comparison between spectroscopic findings (XPS, UPS, and EES) for PPy and model calculations has led some researchers to conclude that 10 is the minimum number of monomeric units in a PPy chain, with the maximum within one order of magnitude [106,109]. By electropolymerizing α,α′-tritium-labeled β,β′-dimethylpyrrole and comparing the tritium activity in the monomer and the polymer, Nazzal and Street obtained molecular weights that indicate chain lengths of between 100 and 1000 Py units [110]. By contrast, the oligomer approach has shown that oligomerization in solution produces soluble oligomers with chain lengths between 8 and 12 units. During solid-state polymerization in dependence on the applied electrode potential, linear chains up to 30–60 units as well as crosslinked networks are generated [39,86].

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IV. CHARgE STORAgE MECHANISM IN CONDUCTINg POLyMERS A.

REDOX PROPERTIES OF OLIGOMERS AND POLYMERS

CPs, provided they are in their neutral state, are initially insulators. Their metal-like properties, that is, their high conductivity and optical reflectivity, become obvious only after doping. As already mentioned, this doping corresponds to an electrochemical oxidation in the case of p-doping and to a reduction in the case of n-doping. Suitable redox reagents are either chemical electron acceptors, such as iodine, or electron donors, such as potassium naphthalide, or the process may be electrochemically induced via an electrochemical cell. Because of the redox reaction, the polymer chain is negatively charged in the case of reduction and positively charged in the case of oxidation. To maintain electroneutrality, the appropriate counterions diffuse into the polymer during charging and out of the polymer during discharging. The most prominent electrochemical technique to monitor such charging processes is cyclic voltammetry. Theoretical concepts that describe the voltammetric response during charging and discharging of redox-active films were developed more than four decades ago [111]. In the ideal case of a simple one-electron transfer, reversible cyclic voltammograms should show completely symmetrical and mirror-image cathodic and anodic waves with identical peak potentials and current levels (Figure 41.8). The current in the reversible case is then i=

n2 F 2 AΓ T v ⋅ expΘ RT (1 + expΘ)2

(41.2)

where Θ = (nF/RT)(E−E 0) ΓT = Γ0 + ΓR

Ep = E 0 i

ip

–0.2

0.0

0.2

(E–E 0) (V)

FIgURE 41.8 Theoretical cyclic voltammogram for a thin-layer film with one-redox center.

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correspond to the total surface covered with reduced and oxidized sites. The other parameters have their usual meanings. Apart from the mirror symmetry of the waves, it is also characteristic that, in contrast to measurements obtained with soluble redox systems, current i and the scan rate v are directly proportional to each other. The earlier statements are valid for monomolecular layers only. In the case of redox-active polymer films with layer thickness up into the µ-range, as are often produced by electropolymerization, account must also be taken of the fact that diffusion sets in. All variants of the diffusion mass transport are possible from the finite limiting case to semi-infinite diffusion, depending on film thickness, the values of the formal diffusion coefficients, and the experimental timescales used. For voltammetric experiments, this implies that as the sweep rate v increases, there must be a shift from mirror symmetrical CV diagrams with i proportional to v to the classic, asymmetrical voltammograms with i proportional v1/2. Although the potentiodynamic charging and discharging of CPs produce voltammograms of very different shape depending on type and polymerization conditions, one frequently finds CV diagrams (Figure 41.9) with a very similar shape. This is particularly the case when polymeric films are produced under mild conditions at low formation potentials [7,14c,23,50,104]. Characteristic features of these systems are, in the case of the p-doping, a steep anodic wave at the start of charging, followed by a broad, flat plateau as potential increases. In the reverse scan, a potential-shifted cathodic wave appears at the negative end of the capacity-like plateau. In order to characterize the charging/discharging processes, two determining factors should be considered. The first one concerns the amount of doping and insertion. The second one is related to the detailed understanding of the redox mechanisms in CPs. A simple strategy to describe the charging or doping level of CPs is the mole fraction of the corresponding monomers that are charged. The optimum doping level for PPy or PTh is about 0.33, but can have very much lower values. This depends, inter alia, on structure and the applied charging potential, but is also influenced by environmental parameters such as the solvent or supporting electrolyte [112–116]. From the respective doping levels, one can deduce that, upon oxidation of PPy, and often on that of PTh as well, formally every third or fourth heterocycle is charged, whereas in the case of PPP, provided that the experiments take place in common solvents such as PC, only every sixth monomeric unit is charged [114]. Moving to a higher doping level, for example, 0.5, a charge is forced in every second monomeric subunit, which, of course, induces a coulombic repulsion. Therefore, to achieve this, a higher electrode potential must be applied, normally resulting in overoxidation effects [117]. Through the overoxidation process, a degradation of the polymer occurs, often induced by

1 μA

–1.0

–0.5 0.0 E (vs. Ag/AgCl) (V)

0.5

FIgURE 41.9 Cyclic voltammogram of the oxidation of poly(ethylenedioxythiophene) in CH2Cl2/0.1 M TBAPF6, T = 273 K, v = 10 mV s−1. (From Dietrich, M., Dissertation, Freiburg, Germany, 1990.)

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16 1° scan 14 2° scan 12 10 8 6 4 2 0 –2 –4 –6 –8 –10 –1.5 –1.0 –0.5 –0.0

1000 100 10 1 0.1 0.01 1E – 3

log (conductance/mS)

I/μA

Electrochemistry of Conducting Polymers

1E – 4 1E – 5 0.5 1.0 1.5 2.0 E (V) vs. Ag/AgCl

2.5

3.0

3.5

FIgURE 41.10 Cyclic voltammetry and in  situ conductance measurements of poly(4,4′-dimethoxybithiophene) in CH2Cl2, 0.1 M TBAPF6, v = 5 mV s−1, T = 273 K. Black thin lines, cyclic voltammograms; black thick lines, conductance as a function of potential measured during potentiodynamic cycling, logarithmic scale representation of conductance. (From Espindola, P., Dissertation, Freiburg, Germany, 2005.)

nucleophilic solvents such as water or nitriles. Therefore, overoxidized materials can be discharged only to a limited extent, which is a serious drawback for applications. In principle, a doping level of 1 should be possible, such that every monomeric unit bears a positive charge after doping. Up to now, only the very stable poly(4,4′-dimethoxy-bithiophene) system [118] has been charged up to this doping level. A characteristic feature of such a perfectly charged system is that the current in a voltammetric experiment drops to zero after passing the highest available redox state (Figure 41.10). This gives evidence that the capacity-like plateau results from a faradaic process. Detailed information about the correlation between redox states and structure was obtained by measurements of charging/discharging properties of well-defined, monodisperse oligo(p-phenylenevinylenes) in solution and under solid-state conditions (Figure 41.11) [45,119]. The analysis of the reduction data showed that the number of accessible redox states increases with increasing chain length of the system resulting in the superposition of redox states over a broad potential range for long chain lengths. Therefore, the controversial capacity-like plateau that generally appears in voltammograms of CPs can be shown to refer to faradaic redox processes. Moreover, the voltammetric signal of the steep anodic wave at the beginning of the charging certainly belongs to a close superposition of several redox states, probably up to a level of tetracations or even more for longer chains. In the case of oligo(p-phenylenevinylene) with six phenylenevinylene units, at least seven redox states can be found in the potential range between −2.0 and 3.0 V (Figure 41.11). The fact that the polymeric material is normally polydisperse additionally favors the superposition of redox states. These findings clarify that the electrochemical charging process of CPs should be described by a sequence of discrete but overlapping redox steps. It implies that a model at the molecular level fits best the redox properties of conjugated materials. A band model as favored by physicists seems to be unlikely (see the following discussion on theoretical models). Equivalent results were obtained with other series of conjugated oligomers [45]. Especially interesting are voltammetric measurements of β-carotenes, which are oligomeric model compounds of PA [57b]. Again, discrete redox steps are visible, the potential gaps between them decreasing with increasing chain length indicating that the redox behavior of highly redox-active oligoenes is in principle similar to that of other chainlike conjugated systems.

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n=3

n=4

0.2 μA

n=5

0.1 μA

n=6 0.01 μA

0.005 μA

n = ? (∞)

–3.5 –3.0 –2.5 –2.0 –1.5 E vs. Ag/AgCl (V)

–3.5 –3.0 –2.5 –2.0 –1.5 –1.0 E vs. Ag /AgCl (V)

FIgURE 41.11 Cyclic voltammograms of the reduction of (left) oligo(p-phenylenevinylenes) in solution (T = −65°C, v = 100 mV s−1) in THF/TBAPF6 and (right) under solid-state conditions (T = −65°C, v between 10 and 50 mV s−1) in DMA/TBABr. (Meerholz, K., Gregorius, H., Müllen, K., and Heinze, J.: Adv. Mat. 1994. 6. 671. Copyright Wiley-VCH Verlag GmbH & Co. KGaA. Reproduced with permission.)

All these data complement experimental and theoretical results already presented in the older literature [120,121]. The following general trends have been established for chainlike conjugated oligomers as a function of increasing chain length [57b,122]: • Redox states of identical charge (e.g., mono- or diion) shift toward lower energies. For long chains, the energies of low redox states gradually approach a common convergence limit. Mono- and diionic redox transitions pairwise degenerate. • Adding successive monomeric subunits in the molecular chain enlarges the number of accessible redox states. • The energy gap widens considerably between the lowest and the highest charged states. • The number of redox states is limited and does not exceed the number of monomeric subunits in a chain. • The chemical stability of charged species related to the same redox states increases, and therefore, the tendency for follow-up reactions decreases [123]. A very important observation within the series of all these oligomers has been that the stabilization of diionic states does not increase relative to monoionic states. The experimental data reveal that

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the generation of higher redox states essentially requires more or at least the same energy than the formation of monoionic states (Figure 41.11). An additional phenomenon in many voltammetric experiments with CPs is the conspicuous separation between the wave of anodic charging and cathodic discharging (hysteresis) (Figures 41.5, 41.7, and 41.9). It was initially interpreted as a kinetic effect of slow heterogeneous charge transfer [7,124] or conformational changes during charging [125]. As the potentials of the