Comprehensive Organometallic Chemistry IV. Volume 1: Fundamentals [1] 9780128202067

Citation preview

COMPREHENSIVE ORGANOMETALLIC CHEMISTRY IV

COMPREHENSIVE ORGANOMETALLIC CHEMISTRY IV EDITORS-IN-CHIEF

GERARD PARKIN Department of Chemistry, Columbia University, New York, NY, United States

KARSTEN MEYER Department of Chemistry and Pharmacy, Friedrich-Alexander-Universität, Erlangen, Germany

DERMOT O’HARE Department of Chemistry, University of Oxford, Oxford, United Kingdom

VOLUME 1

FUNDAMENTALS VOLUME EDITOR

PATRICK L. HOLLAND Yale University, New Haven, United States

Elsevier Radarweg 29, PO Box 211, 1000 AE Amsterdam, Netherlands The Boulevard, Langford Lane, Kidlington, Oxford OX5 1GB, United Kingdom 50 Hampshire Street, 5th Floor, Cambridge, MA 02139, United States Copyright © 2022 Elsevier Ltd. All rights reserved No part of this publication may be reproduced or transmitted in any form or by any means, electronic or mechanical, including photocopying, recording, or any information storage and retrieval system, without permission in writing from the publisher. Details on how to seek permission, further information about the Publisher’s permissions policies and our arrangements with organizations such as the Copyright Clearance Center and the Copyright Licensing Agency, can be found at our website: www.elsevier.com/permissions. This book and the individual contributions contained in it are protected under copyright by the Publisher (other than as may be noted herein). Notices Knowledge and best practice in this field are constantly changing. As new research and experience broaden our understanding, changes in research methods, professional practices, or medical treatment may become necessary. Practitioners and researchers may always rely on their own experience and knowledge in evaluating and using any information, methods, compounds, or experiments described herein. In using such information or methods they should be mindful of their own safety and the safety of others, including parties for whom they have a professional responsibility. To the fullest extent of the law, neither the Publisher nor the authors, contributors, or editors, assume any liability for any injury and/or damage to persons or property as a matter of products liability, negligence or otherwise, or from any use or operation of any methods, products, instructions, or ideas contained in the material herein.

Library of Congress Cataloging-in-Publication Data A catalog record for this book is available from the Library of Congress British Library Cataloguing-in-Publication Data A catalogue record for this book is available from the British Library ISBN 978-0-12-820206-7 For information on all publications visit our website at http://store.elsevier.com

Publisher: Oliver Walter Acquisition Editor: Blerina Osmanaj Content Project Manager: Claire Byrne Associate Content Project Manager: Fahmida Sultana Designer: Christian Bilbow

CONTENTS OF VOLUME 1 Editor Biographies

vii

Contributors to Volume 1

xiii

Preface 1.01

xv

Introduction: Volume I

1

Patrick L Holland

1.02

Models for Understanding Main Group and Transition Metal Bonding

2

Aaron L Odom

1.03

Reversible Homolysis of Metal-Carbon Bonds

31

Maxime Michelas, Christophe Fliedel, and Rinaldo Poli

1.04

Very Low Oxidation States in Organometallic Chemistry

86

C Gunnar Werncke

1.05

Very High Oxidation States in Organometallic Chemistry

109

Moritz Malischewski

1.06

Characterization Methods for Paramagnetic Organometallic Complexes

135

Aleksa Radovic, Shilpa Bhatia, and Michael L Neidig

1.07

Computational Methods in Organometallic Chemistry

176

S Chantal E Stieber

1.08

f-Element Organometallic Single-Molecule Magnets

211

Richard A Layfield, Christopher GT Price, and Siobhan R Temple

1.09

Electrochemistry in Organometallic Chemistry

249

Julie A Hopkins Leseberg, Wade C Henke, and James D Blakemore

1.10

Organometallic Photosensitizers

284

Thomas S Teets and Yanyu Wu

1.11

Organometallic Chemistry of NHCs and Analogues

339

Liang Deng and Zhenbo Mo

1.12

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

373

Anton W Tomich, Varun Tej, Sergio Lovera, Isaac Banda, Steven Fisher, Matthew Asay, and Vincent Lavallo

1.13

Redox-Active Ligands in Organometallic Chemistry

421

Errikos Kounalis and Daniël LJ Broere

1.14

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

442

Elizabeth T Papish, Sanjit Das, Weerachai Silprakob, Chance M Boudreaux, and Sonya Manafe

1.15

Introduction to the Organometallic Chemistry of Carbon Dioxide

474

Charles W Machan

1.16

Alkane s-Complexes

502

Rowan D Young

v

vi

1.17

Contents of Volume 1

Dinitrogen Binding and Functionalization

521

Jeremy E Weber, Samuel M Bhutto, Alexandre T-Y Genoux, and Patrick L Holland

1.18

Lewis Acid Participation in Organometallic Chemistry

555

Julia B Curley, Nilay Hazari, and Tanya M Townsend

1.19

Organometallic Chemistry on Oxide Surfaces

583

Matthew P Conley, Jiaxin Gao, Winn Huynh, Jessica Rodriguez, and Kavyasripriya K Samudrala

1.20

Separation Strategies in Organometallic Catalysis

609

Fernanda G Mendonça and R Tom Baker

1.21

Impurities in Organometallic Catalysis Nicholas E Leadbeater

635

EDITOR BIOGRAPHIES Editors in Chief Karsten Meyer studied chemistry at the Ruhr University Bochum and performed his Ph.D. thesis work on the molecular and electronic structure of first-row transition metal complexes under the direction of Professor Karl Wieghardt at the Max Planck Institute in Mülheim/Ruhr (Germany). He then proceeded to gain research experience in the laboratory of Professor Christopher Cummins at the Massachusetts Institute of Technology (USA), where he appreciated the art of synthesis and developed his passion for the coordination chemistry and reactivity of uranium complexes. In 2001, he was appointed to the University of California, San Diego, as an assistant professor and was named an Alfred P. Sloan Fellow in 2004. In 2006, he accepted an offer (C4/W3) to be the chair of the Institute of Inorganic & General Chemistry at the Friedrich-Alexander-University ErlangenNürnberg (FAU), Germany. Among his awards and honors, he was elected a lifetime honorary member of the Israel Chemical Society and a fellow of the Royal Society of Chemistry (UK). Karsten received the Elhuyar-Goldschmidt Award from the Royal Society of Chemistry of Spain, the Ludwig Mond Award from the RSC (UK), and the Chugaev Commemorative Medal from the Russian Academy of Sciences. He has also enjoyed visiting professorship positions at the universities of Manchester (UK) and Toulouse (F) as well as the Nagoya Institute of Technology (JP) and ETH Zürich (CH). The Meyer lab research focuses on the synthesis of custom-tailored ligand environments and their transition and actinide metal coordination complexes. These complexes often exhibit unprecedented coordination modes, unusual electronic structures, and, consequently, enhanced reactivities toward small molecules of biological and industrial importance. Interestingly, Karsten’s favorite molecule is one that exhibits little reactivity: the Th symmetric U(dbabh)6. Dermot O’Hare was born in Newry, Co Down. He studied at Balliol College, Oxford University, where he obtained his B.A., M.A., and D.Phil. degrees under the direction of Professor M.L.H. Green. In 1985, he was awarded a Royal Commission of 1851 Research Fellowship, during this Fellowship he was a visiting research fellow at the DuPont Central Research Department, Wilmington, Delaware in 1986–87 in the group led by Prof. J.S. Miller working on molecular-based magnetic materials. In 1987 he returned to Oxford to a short-term university lectureship and in 1990 he was appointed to a permanent university position and a Septcentenary Tutorial Fellowship at Balliol College. He has previously been honored by the Institüt de France, Académie des Sciences as a leading scientist in Europe under 40 years. He is currently professor of organometallic and materials chemistry in the Department of Chemistry at the University of Oxford. In addition, he is currently the director of the SCG-Oxford Centre of Excellence for chemistry and associate head for business & innovation in the Mathematics, Physical and Life Sciences Division. He leads a multidisciplinary research team that works across broad areas of catalysis and nanomaterials. His research is specifically targeted at finding solutions to global issues relating to energy, zero carbon, and the circular economy. He has been awarded numerous awards and prizes for his creative and

vii

viii

Editor Biographies

ground-breaking work in inorganic chemistry, including the Royal Society Chemistry’s Sir Edward Frankland Fellowship, Ludwig Mond Prize, Tilden Medal, and Academia–Industry Prize and the Exxon European Chemical and Engineering Prize. Gerard Parkin received his B.A., M.A., and D.Phil. degrees from the Queen’s College, Oxford University, where he carried out research under the guidance of Professor Malcolm L.H. Green. In 1985, he moved to the California Institute of Technology as a NATO postdoctoral fellow to work with Professor John E. Bercaw. He joined the Faculty of Columbia University as assistant professor in 1988 and was promoted to associate professor in 1991 and to professor in 1994. He served as chairman of the Department from 1999 to 2002. He has also served as chair of the New York Section of the American Chemical Society, chair of the Inorganic Chemistry and Catalytic Science Section of the New York Academy of Sciences, chair of the Organometallic Subdivision of the American Chemical Society Division of Inorganic Chemistry, and chair of the Gordon Research Conference in Organometallic Chemistry. He is an elected fellow of the American Chemical Society, the Royal Society of Chemistry, and the American Association for the Advancement of Science, and is the recipient of a variety of international awards, including the ACS Award in pure chemistry, the ACS Award in organometallic chemistry, the RSC Corday Morgan Medal, the RSC Award in organometallic chemistry, the RSC Ludwig Mond Award, and the RSC Chem Soc Rev Lecture Award. He is also the recipient of the United States Presidential Award for Excellence in Science, Mathematics and Engineering Mentoring, the United States Presidential Faculty Fellowship Award, the James Flack Norris Award for Outstanding Achievement in the Teaching of Chemistry, the Columbia University Presidential Award for Outstanding Teaching, and the Lenfest Distinguished Columbia Faculty Award. His principal research interests are in the areas of synthetic, structural, and mechanistic inorganic chemistry.

Volume Editors Simon Aldridge is professor of chemistry at the University of Oxford and director of the UKRI Centre for Doctoral Training in inorganic chemistry for Future Manufacturing. Originally from Shrewsbury, England, he received both his B.A. and D.Phil. degrees from the University of Oxford, the latter in 1996 for work on hydride chemistry under the supervision of Tony Downs. After post-doctoral work as a Fulbright Scholar at Notre Dame with Tom Fehlner, and at Imperial College London (with Mike Mingos), he took up his first academic position at Cardiff University in 1998. He returned to Oxford in 2007, being promoted to full professor in 2010. Prof. Aldridge has published more than 230 papers to date and is a past winner of the Dalton Transactions European Lectureship (2009), the Royal Society of Chemistry’s Main Group Chemistry (2010) and Frankland Awards (2018), and the Forschungspreis of the Alexander von Humboldt Foundation (2021). Prof. Aldridge’s research interests are primarily focused on main group organometallic chemistry, and in particular the development of compounds with unusual electronic structure, and their applications in small molecule activation and catalysis (website: http:// aldridge.web.ox.ac.uk). (Picture credit: John Cairns)

Editor Biographies

ix

Eszter Boros is associate professor of chemistry at Stony Brook University with courtesy appointments in radiology and pharmacology at Stony Brook Medicine. Eszter obtained her M.Sc. (2007) at the University of Zurich, Switzerland and her Ph.D. (2011) in chemistry from the University of British Columbia, Canada. She was a postdoc (2011–15) and later instructor (2015–17) in radiology at Massachusetts General Hospital and Harvard Medical School. In 2017, Eszter was appointed as assistant professor of chemistry at Stony Brook University, where her research group develops new approaches to metal-based diagnostics and therapeutics at the interfaces of radiochemistry, inorganic chemistry and medicine. Her lab’s work has been extensively recognized; Eszter holds various major federal grants (NSF CAREER Award, NIH NIBIB R21 Trailblazer, NIH NIGMS R35 MIRA) and has been named a 2020 Moore Inventor Fellow, the 2020 Jonathan L. Sessler Fellow (American Chemical Society, Inorganic Division), recipient of a 2021 ACS Infectious Diseases/ACS Division of Biological Chemistry Young Investigator Award (American Chemical Society), and was also named a 2022 Alfred P. Sloan Research Fellow in chemistry. Scott R. Daly is associate professor of chemistry at the University of Iowa in the United States. After spending 3 years in the U.S. Army, he obtained his B.S. degree in chemistry in 2006 from North Central College, a small liberal arts college in Naperville, Illinois. He then went on to receive his Ph.D. at the University of Illinois at Urbana-Champaign in 2010 under the guidance of Professor Gregory S. Girolami. His thesis research focused on the synthesis and characterization of chelating borohydride ligands and their use in the preparation of volatile metal complexes for chemical vapor deposition applications. In 2010, he began working as a Seaborg postdoctoral fellow with Drs. Stosh A. Kozimor and David L. Clark at Los Alamos National Laboratory in Los Alamos, New Mexico. His research there concentrated on the development of ligand K-edge X-ray absorption spectroscopy (XAS) to investigate covalent metal–ligand bonding and electronic structure variations in actinide, lanthanide, and transition metal complexes with metal extractants. He started his independent career in 2012 at George Washington University in Washington, DC, and moved to the University of Iowa shortly thereafter in 2014. His current research interests focus on synthetic coordination chemistry and ligand design with emphasis on the development of chemical and redox noninnocent ligands, mechanochemical synthesis and separation methods, and ligand K-edge XAS. His research and outreach efforts have been recognized with an Outstanding Faculty/Staff Advocate Award from the University of Iowa Veterans Association (2016), a National Science Foundation CAREER Award (2017), and a Hawkeye Distinguished Veterans Award (2018). He was promoted to associate professor with distinction as a College of Liberal Arts and Sciences Deans Scholar in 2020. Lena J. Daumann is currently professor of bioinorganic and coordination chemistry at the Ludwig Maximilian Universität in Munich. She studied chemistry at the University of Heidelberg working with Prof. Peter Comba and subsequently conducted her Ph.D. at the University of Queensland (Australia) from 2010 to 2013 holding IPRS and UQ Centennial fellowships. In 2013 she was part of the Australian Delegation for the 63rd Lindau Nobel Laureate meeting in chemistry. Following postdoctoral stays at UC Berkeley with Prof. Ken Raymond (2013–15) and in Heidelberg, funded by the Alexander von Humboldt Foundation, she started her independent career at the LMU Munich in 2016. Her bioinorganic research group works on elucidating the role of lanthanides for bacteria as well as on iron enzymes and small biomimetic complexes that play a role in epigenetics and DNA repair. Daumann’s teaching and research have been recognized with numerous awards and grants. Among them are the national Ars Legendi Prize for chemistry and the Therese von Bayern Prize in 2019 and the Dozentenpreis of the “Fonds der Chemischen Industrie“ in 2021. In 2018 she was selected as fellow for the Klaus Tschira Boost Fund by the German Scholars Organisation and in 2020 she received a Starting grant of the European Research Council to study the uptake of lanthanides by bacteria.

x

Editor Biographies

Derek P. Gates hails from Halifax, Nova Scotia (Canada) where he completed his B.Sc. (Honours Chemistry) degree at Dalhousie University in 1993. He completed his Ph.D. degree under the supervision of Professor Ian Manners at the University of Toronto in 1997. He then joined the group of Professor Maurice Brookhart as an NSERC postdoctoral fellow at the University of North Carolina at Chapel Hill (USA). He began his independent research career in 1999 as an assistant professor at the University of British Columbia in Vancouver (Canada). He has been promoted through the ranks and has held the position of professor of chemistry since 2011. At UBC, he has received the Science Undergraduate Society—Teaching Excellence Award, the Canadian National Committee for IUPAC Award, and the Chemical Society of Canada—Strem Chemicals Award for pure or applied inorganic chemistry. His research interests bridge the traditional fields of inorganic and polymer chemistry with particular focus on phosphorus chemistry. Key topics include the discovery of novel structures, unusual bonding, new reactivity, along with applications in catalysis and materials science. Patrick Holland performed his Ph.D. research in organometallic chemistry at UC Berkeley with Richard Andersen and Robert Bergman. He then learned about bioinorganic chemistry through postdoctoral research on copper-O2 and copper-thiolate chemistry with William Tolman at the University of Minnesota. His independent research at the University of Rochester initially focused on systematic development of the properties and reactions of three-coordinate complexes of iron and cobalt, which can engage in a range of bond activation reactions and organometallic transformations. Since then, his research group has broadened its studies to iron-N2 chemistry, reactive metal–ligand multiple bonds, iron–sulfur clusters, engineered metalloproteins, redox-active ligands, and solar fuel production. In 2013, Prof. Holland moved to Yale University, where he is now Conkey P. Whitehead Professor of Chemistry. His research has been recognized with an NSF CAREER Award, a Sloan Research Award, Fulbright and Humboldt Fellowships, a Blavatnik Award for Young Scientists, and was elected as fellow of the American Association for the Advancement of Science. In the area of N2 reduction, his group has established molecular principles to weaken and break the strong N–N bond, in order to use this abundant resource for energy and synthesis. His group has made a particular effort to gain an insight into iron chemistry relevant to nitrogenase, the enzyme that reduces N2 in nature. His group also maintains an active program in the use of inexpensive metals for transformations of alkenes. Mechanistic details are a central motivation to Prof. Holland and the wonderful group of over 80 students with whom he has worked. Steve Liddle was born in Sunderland in the North East of England and gained his B.Sc. (Hons) and Ph.D. from Newcastle University. After postdoctoral fellowships at Edinburgh, Newcastle, and Nottingham Universities he began his independent career at Nottingham University in 2007 with a Royal Society University Research Fellowship. This was held in conjunction with a proleptic Lectureship and he was promoted through the ranks to associate professor and reader in 2010 and professor of inorganic chemistry in 2013. He remained at Nottingham until 2015 when he was appointed professor and head of inorganic chemistry and co-director of the Centre for Radiochemistry Research at The University of Manchester. He has been a recipient of an EPSRC Established Career Fellowship and ERC Starter and Consolidator grants. He is an elected fellow of The Royal Society of Edinburgh and fellow of the Royal Society of Chemistry and he is vice president to the Executive Committee of the European Rare Earth and Actinide Society. His principal research interests are focused on f-element chemistry, involving exploratory synthetic chemistry coupled to detailed electronic structure and reactivity studies to elucidate structure-bonding-property relationships. He is the recipient of a variety of prizes, including the IChemE Petronas Team Award for Excellence in Education and Training, the RSC Sir Edward Frankland Fellowship, the RSC Radiochemistry

Editor Biographies

xi

Group Bill Newton Award, a 41st ICCC Rising Star Award, the RSC Corday-Morgan Prize, an Alexander von Humboldt Foundation Friedrich Wilhelm Bessel Research Award, the RSC Tilden Prize, and an RSC Dalton Division Horizon Team Prize. He has published over 220 research articles, reviews, and book chapters to date. David Liptrot received his MChem (Hons) in chemistry with Industrial Training from the University of Bath in 2011 and remained there to undertake a Ph.D. on group 2 catalysis in the laboratory of Professor Mike Hill. After completing this in 2015 he took up a Lindemann Postdoctoral Fellowship with Professor Philip Power FRS (University of California, Davis, USA). In 2017 he began his independent career returning to the University of Bath and in 2019 was awarded a Royal Society University Research Fellowship. His interests concern new synthetic methodologies to introduce main group elements into functional molecules and materials.

David P. Mills hails from Llanbradach and Caerphilly in the South Wales Valleys. He completed his MChem (2004) and Ph.D. (2008) degrees at Cardiff University, with his doctorate in low oxidation state gallium chemistry supervised by Professor Cameron Jones. He moved to the University of Nottingham in 2008 to work with Professor Stephen Liddle for postdoctoral studies in lanthanide and actinide methanediide chemistry. In 2012 he moved to the University of Manchester to start his independent career as a lecturer, where he has since been promoted to full professor of inorganic chemistry in 2021. Although he is interested in all aspects of nonaqueous synthetic chemistry his research interests are currently focused on the synthesis and characterization of f-block complexes with unusual geometries and bonding regimes, with the aim of enhancing physicochemical properties. He has been recognized for his contributions to both research and teaching with prizes and awards, including a Harrison-Meldola Memorial Prize (2018), the Radiochemistry Group Bill Newton Award (2019), and a Team Member of the Molecular Magnetism Group for the Dalton Division Horizon Prize (2021) from the Royal Society of Chemistry. He was a Blavatnik Awards for Young Scientists in the United Kingdom Finalist in Chemistry in 2021 and he currently holds a European Research Council Consolidator Grant. Ian Tonks is the Lloyd H. Reyerson professor at the University of MinnesotaTwin Cities, and associate editor for the ACS journal Organometallics. He received his B.A. in chemistry from Columbia University in 2006 and performed undergraduate research with Prof. Ged Parkin. He earned his Ph.D. in 2012 from the California Institute of Technology, where he worked with Prof. John Bercaw on olefin polymerization catalysis and early transition metal-ligand multiply bonded complexes. After postdoctoral research with Prof. Clark Landis at the University of Wisconsin, Madison, he began his independent career at the University of Minnesota in 2013 and earned tenure in 2019. His current research interests are focused on the development of earth abundant, sustainable catalytic methods using early transition metals, and also on catalytic strategies for incorporation of CO2 into polymers. Prof. Tonks’ work has recently been recognized with an Outstanding New Investigator Award from the National Institutes of Health, an Alfred P. Sloan Fellowship, a Department of Energy CAREER award, and the ACS Organometallics Distinguished Author Award, among others. Additionally, Prof. Tonks’ service toward improving academic safety culture was recently recognized with the 2021 ACS Division of Chemical Health and Safety Graduate Faculty Safety Award.

xii

Editor Biographies

Timothy H. Warren is the Rosenberg professor and chairperson in the Department of Chemistry at Michigan State University. He obtained his B.S. from the University of Illinois at Urbana-Champaign in 1992 and Ph.D. from the Massachusetts Institute of Technology in 1997. After 2 years of postdoctoral research at the Organic Chemistry Institute of the University of Münster, Germany with Prof. Dr. Gerhart Erker, Dr. Warren started his independent career at Georgetown University in 1999 where he was named the Richard D. Vorisek professor of chemistry in 2014. He moved to Michigan State University in 2021. Prof. Warren’s research interests span synthetic and mechanistic inorganic, organometallic, and bioinorganic chemistry with a focus on catalysis. His research group develops environmentally friendly methods for organic synthesis via C–H functionalization, explores the interconversion of nitrogen and ammonia as carbon-free fuels, and decodes ways that biology communicates using nitric oxide as a molecular messenger. Mechanistic studies on these chemical reactions catalyzed by metal ions such as iron, nickel, copper, and zinc enable new insights for the development of useful catalysts for synthesis and energy applications as well as lay the mechanistic groundwork to understand biochemical nitric oxide misregulation. Dr. Warren received the NSF CAREER Award, chaired the 2019 Inorganic Reaction Mechanisms Gordon Research Conference, and has served on the ACS Division of Inorganic Chemistry executive board and on the editorial boards of Inorganic Synthesis, Inorganic Chemistry, and Chemical Society Reviews.

CONTRIBUTORS TO VOLUME 1 Matthew Asay Universal Display Corporation, Ewing, NJ, United States R Tom Baker Department of Chemistry and Biomolecular Sciences and Centre for Catalysis Research and Innovation, University of Ottawa, Ottawa, ON, Canada Isaac Banda University of California Riverside, Riverside, CA, United States Shilpa Bhatia Department of Chemistry, University of Rochester, Rochester, NY, United States Samuel M Bhutto Department of Chemistry, Yale University, New Haven, CT, United States James D Blakemore Department of Chemistry, University of Kansas, Lawrence, KS, United States Chance M Boudreaux Department of Chemistry, The University of Alabama, Tuscaloosa, AL, United States Daniël LJ Broere Organic Chemistry and Catalysis, Debye Institute for Nanomaterials Science Faculty of Science, Utrecht University, Universiteitsweg 99, Utrecht, The Netherlands Matthew P Conley Department of Chemistry, University of California, Riverside, CA, United States Julia B Curley Department of Chemistry, Yale University, New Haven, CT, United States Sanjit Das Department of Chemistry, The University of Alabama, Tuscaloosa, AL, United States

Liang Deng State Key Laboratory of Organometallic Chemistry, Shanghai Institute of Organic Chemistry, Chinese Academy of Sciences, Shanghai, PR China Steven Fisher Lawrence Livermore National Laboratory, Livermore, CA, United States Christophe Fliedel Laboratoire de Chimie de Coordination, UPR CNRS 8241, Toulouse, France Jiaxin Gao Department of Chemistry, University of California, Riverside, CA, United States Alexandre T-Y Genoux Department of Chemistry, Yale University, New Haven, CT, United States Nilay Hazari Department of Chemistry, Yale University, New Haven, CT, United States Wade C Henke Department of Chemistry, University of Kansas, Lawrence, KS, United States Patrick L Holland Department of Chemistry, Yale University, New Haven, CT, United States Julie A Hopkins Leseberg Department of Chemistry, University of Kansas, Lawrence, KS, United States Winn Huynh Department of Chemistry, University of California, Riverside, CA, United States Errikos Kounalis Organic Chemistry and Catalysis, Debye Institute for Nanomaterials Science Faculty of Science, Utrecht University, Universiteitsweg 99, Utrecht, The Netherlands

xiii

xiv

Contributors to Volume 1

Vincent Lavallo University of California Riverside, Riverside, CA, United States

Christopher GT Price Department of Chemistry, School of Life Sciences, University of Sussex, Brighton, United Kingdom

Richard A Layfield Department of Chemistry, School of Life Sciences, University of Sussex, Brighton, United Kingdom

Aleksa Radovic Department of Chemistry, University of Rochester, Rochester, NY, United States

Nicholas E Leadbeater Department of Chemistry, University of Connecticut, Storrs, CT, United States

Jessica Rodriguez Department of Chemistry, University of California, Riverside, CA, United States

Sergio Lovera University of California Riverside, Riverside, CA, United States

Kavyasripriya K Samudrala Department of Chemistry, University of California, Riverside, CA, United States

Charles W Machan Department of Chemistry, University of Virginia, Charlottesville, VA, United States

Weerachai Silprakob Department of Chemistry, The University of Alabama, Tuscaloosa, AL, United States

Moritz Malischewski Freie Universität Berlin, Institut für Chemie und Biochemie—Anorganische Chemie, Berlin, Germany

S Chantal E Stieber Department of Chemistry & Biochemistry, California State Polytechnic University, Pomona, CA, United States

Sonya Manafe Department of Chemistry, The University of Alabama, Tuscaloosa, AL, United States Fernanda G Mendonça Department of Chemistry and Biomolecular Sciences and Centre for Catalysis Research and Innovation, University of Ottawa, Ottawa, ON, Canada Maxime Michelas Laboratoire de Chimie de Coordination, UPR CNRS 8241, Toulouse, France Zhenbo Mo State Key Laboratory and Institute of Elemento-Organic Chemistry, College of Chemistry, Nankai University, Tianjin, PR China Michael L Neidig Department of Chemistry, University of Rochester, Rochester, NY, United States Aaron L Odom Department of Chemistry, Michigan State University, East Lansing, MI, United States Elizabeth T Papish Department of Chemistry, The University of Alabama, Tuscaloosa, AL, United States Rinaldo Poli Laboratoire de Chimie de Coordination, UPR CNRS 8241, Toulouse, France

Thomas S Teets Department of Chemistry, University of Houston, Houston, TX, United States Varun Tej University of California Riverside, Riverside, CA, United States Siobhan R Temple Department of Chemistry, School of Life Sciences, University of Sussex, Brighton, United Kingdom Anton W Tomich University of California Riverside, Riverside, CA, United States Tanya M Townsend Department of Chemistry, Yale University, New Haven, CT, United States Jeremy E Weber Department of Chemistry, Yale University, New Haven, CT, United States C Gunnar Werncke Chemistry Department, Philipps-University Marburg, Marburg, Germany Yanyu Wu Department of Chemistry, University of Houston, Houston, TX, United States Rowan D Young National University of Singapore, Singapore, Singapore

PREFACE Published 40 years ago in 1982, the first edition of Comprehensive Organometallic Chemistry (COMC) provided an invaluable resource that enabled chemists to become efficiently informed of the properties and reactions of organometallic compounds of both the main group and transition metals. This area of chemistry continued to develop at a rapid pace such that it necessitated the publication of subsequent editions, namely Comprehensive Organometallic Chemistry II (COMC2) in 1995 and Comprehensive Organometallic Chemistry III (COMC3) in 2007. Organometallic chemistry has continued to be vibrant in the 15 years following the publication of COMC3, not only by affording compounds with novel structures and reactivity but also by having important applications in organic syntheses and industrial processes, as illustrated by the awarding of the 2010 Nobel prize to Heck, Negishi, and Suzuki for the development of palladium-catalyzed cross couplings in organic syntheses. Comprehensive Organometallic Chemistry IV (COMC4) thus serves the same important role as its predecessors by providing an indispensable means for researchers and educators to obtain efficiently an up-to-date analysis of a particular aspect of organometallic chemistry. COMC4 comprises 15 volumes, of which the first provides a review of topics concerned with techniques and concepts that feature prominently in current organometallic chemistry, while 5 volumes are devoted to applications that include organic synthesis, materials science, bio-organometallics, metallo-therapy, metallodiagnostics, medicine, and environmental chemistry. In this regard, we are very grateful to the volume editors for their diligent efforts, and the authors for producing high-quality chapters, all of which were written during the COVID-19 pandemic. Finally, we wish to thank the many staff at Elsevier for their efforts to ensure that the project, initiated in the winter of 2018, remained on schedule. Karsten Meyer, Erlangen, March 2022 Dermot O’Hare, Oxford, March 2022 Gerard Parkin, New York, March 2022

xv

1.01

Introduction: Volume I

Patrick L Holland, Department of Chemistry, Yale University, New Haven, CT, United States © 2022 Elsevier Ltd. All rights reserved.

The field of organometallic chemistry has undergone substantial evolution since the publication of the original Comprehensive Organometallic Chemistry in 1982, and even since the appearance of Comprehensive Organometallic Chemistry III in 2007. This has encouraged us to compile a fresh version of Volume 1 that introduces the reader to cross-cutting concepts that are involved in the subsequent volumes. The topics have been chosen to minimize overlap with the chapters in Volume I of Comprehensive Organometallic Chemistry III, and to highlight emerging areas. For example, valence bond theory models for understanding bonding have been used increasingly due to their intuitive content (Odom), and the popularization of redox-active ligands (Broere) has encouraged organometallic chemists to look beyond the formal oxidation state. Chemists have continued to push the boundaries of oxidation states with highly oxidized (Malischewski) and highly reduced (Werncke) complexes, which display amazing reactivity as well. In organometallic chemistry, chemists also strive toward the binding and activation of weak substrates like alkanes (Young), carbon dioxide (Machan), and dinitrogen (Holland). Growing attention to the use of abundant first-row metals has fed a renaissance of paramagnetic organometallic complexes and advanced spectroscopic techniques, which are covered in chapters on computations (Stieber) and on spectroscopy (Neidig). Organometallic chemistry has remained largely driven by ligand design, and a number of chapters explore popular ligand types such as N-heterocyclic carbenes (Deng) and ligands with charge separated from the metal (Lavallo). In addition, influences from outside the coordination sphere are increasingly utilized, and accordingly we include chapters on organometallic systems that incorporate Lewis acid participation (Hazari), proton-responsive ligands (Papish), and attachment to oxide surfaces (Conley). The use of organometallic complexes has also expanded in non-traditional applications like electrochemistry (Blakemore), single-molecule magnets (Layfield), photosensitizers (Teets), and radical reactions (Poli). Finally, practical considerations in catalysis have fueled research into strategies for heterogenization and separation (Baker), and on principles that help the chemist to recognize spurious results from catalyst impurities (Leadbeater). Construction of these extensive reviews requires an immense amount of effort, and all organometallic chemists owe a debt of gratitude to these generous authors for the time they devoted to teaching us about modern aspects of organometallic chemistry. They have all thoughtfully revised to incorporate my suggestions for clarification, and I am confident that the resulting chapters give insights and trends that will guide the next generation of organometallic chemists to even greater heights.

Comprehensive Organometallic Chemistry IV

https://doi.org/10.1016/B978-0-12-820206-7.00166-9

1

1.02

Models for Understanding Main Group and Transition Metal Bonding

Aaron L Odom, Department of Chemistry, Michigan State University, East Lansing, MI, United States © 2022 Elsevier Ltd. All rights reserved.

1.02.1 Introduction 1.02.2 Primogenic repulsion: Orbital size effects in bonding in the periodic table 1.02.3 The 2-center 2-electron heterocovalent bond and polar covalence theory 1.02.4 The 2-center 2-electron dative bond 1.02.5 Complementarity of qualitative hybridization theory and molecular orbital theory 1.02.6 Lone pair bond weakening and its influence on organometallic chemistry 1.02.7 Beyond the 2-center 2-electron bond 1.02.8 Examples of using Bent’s rule, hybridization, and structure in the main group 1.02.9 Hybridization theory and the transition metals 1.02.10 Practical evaluation of metal-ligand bond interactions for applications in catalysis 1.02.11 Concluding remarks Acknowledgments References

1.02.1

2 2 5 12 14 18 20 23 25 27 29 29 29

Introduction

Structure is a central concept that defines and empowers function in biology, chemistry, and physics. Pauling said, “The whole problem of understanding science is, I believe, the problem of relating facts to the concept of structure, first in terms of atoms and then in terms of something still smaller, such as nucleons . . . It is structure that we look for whenever we try to understand anything. All science is built upon this search: we investigate how the cell is built of reticular material, cytoplasm, chromosomes; how crystals aggregate; how atoms are fastened together; how electrons constitute a chemical bond between atoms. We like to understand, and to explain, observed facts in terms of structure.”1 Structure for beginners in chemistry usually consists of viewing molecules with a “hub and strut” model, where the atoms act as a hub in the structure connected by bonding struts represented by lines. Upon this basis, we add layers of complexity depending on the types of atoms and bonds, but molecular structure remains the basis of our models for understanding electronic structure, reactivity, and properties. The atomic “hubs,” from a chemical perspective, are often quite adjustable, and the same element may vary dramatically in the number and types of bonds from one compound to the next. What chemists mean by a line between atoms is also wonderfully diverse, rich, and complex as well. Because a full quantum-mechanical treatment is complex and difficult to conceptualize, organometallic chemists seek simpler models that are more intuitive, yet grounded in quantum mechanics. These simpler expressions of nature may have limitations, but as Roald Hoffmann said, “Chemistry is not mathematics, and even if it bothers some people like hell, there are no theorems of chemistry. Chemical arguments are not falsified by an exception.”2 If one approximation for understanding the bonding fails, we attempt to use a better (sometimes just as simple, sometimes more complex) approximation to understand the system. This chapter is a perspective on “what the lines mean” in some chemical structures. It focuses on relatively simple models for understanding bonding that are readily applied, as these more often help one design the next experiment. This seems fitting, as said in a recent review by Schwerdtfeger, Frenking, and coworkers (emphasis theirs), “Chemical bonding models are not right or wrong, they are more or less useful.”3 In addition, some fundamental principles will be covered that are beneficial for understanding bonding throughout the periodic table. There will be little discussion of historical context, basics of molecular orbital theory, and other topics that are common to undergraduate texts. What is covered will be no less fundamental, but the attempt is made to boil down the discussion to a few (largely pen and paper) models the author finds particularly helpful; these are sometimes supplemented with quantum chemical results, e.g., density functional theory (DFT) and natural bond orbital (NBO) calculations.

1.02.2

Primogenic repulsion: Orbital size effects in bonding in the periodic table

We start with the task of explaining the “primogenic effect,” which will clarify: why the 1st row is so different from the rest of the p-block; why the 1st row transition metals show different behavior (e.g., weaker bonds and weaker ligand fields) than the 2nd and 3rd row; and why the lanthanides show little f-orbital participation in bonding while the actinides show significantly more.

2

Comprehensive Organometallic Chemistry IV

https://doi.org/10.1016/B978-0-12-820206-7.00100-1

Models for Understanding Main Group and Transition Metal Bonding

3

Before describing this effect, we need to review three simple ideas: (1) All the atomic orbitals of an atom are orthogonal to one another. This orthogonality is maintained, not through spatial separation, but through cancelation of bonding and antibonding interactions. The net overlap of any atomic orbital with another is zero due to this cancelation. (2) This orthogonality is maintained through a combination of angular and radial nodes. The 1s orbital, for example, maintains orthogonality with the 2s orbital because the 2s orbital has a node, a change of sign in the wave, so that the two orbitals have no net overlap. (More on this to follow.) For angular differences, the cancelation is readily seen in say the overlap of a p orbital with an s on the same atom (Fig. 1). The constructive overlap, unshaded with unshaded, is exactly canceled by the destructive overlap, unshaded with shaded, which is true of any combination of orbitals with different angular momentum on the same atom. (3) Finally, it is useful to have an expectation for the size of orbitals of different angular momentum with the same principal quantum number (n). For example, when n ¼ 4 there are 4s, 4p, 4d, and 4f orbitals. Which orbital do we expect to be the smallest and which the largest in radius? The angular momentum quantum number l corresponds to the shape of the orbital, with shapes having greater curvature being associated with higher angular momentum. Thinking from a classical perspective for a moment, we can model the electron’s motion around the nucleus as a ball (electron) rotating on an elastic string (Coulombic attraction to the nucleus). If we increase the angular momentum, the elastic string should stretch further from the central point. In other words, we expect that the 4s orbital, which has a lowest angular momentum, would be most compact (closest to nucleus) while the 4f orbital, which has the highest angular momentum, would be the most diffuse (farthest from nucleus). When comparing filled orbitals with the same principal quantum number, lower angular momentum orbitals are expected to have smaller radii. The term “primogenic repulsion” was first published by Pyykkö in 1979,4 and he attributed the phrase to a colleague in the Department of English. The primogenic effect refers to the ability of orbitals of a specific angular momentum appearing for the first time in the Aufbau sequence, i.e., 1s, 2p, 3d, and 4f, to shrink towards the nucleus more than might be expected, simply because there are no orbitals of the same angular momentum in the core below them. All orbitals in a particular shell are orthogonal by their shape due to the l and ml quantum number differences (vide supra). By this principle, most orbitals then have their size set, to some extent, by the necessity of being orthogonal to the orbitals of the same angular momentum below them. In other words, the size of the 1s-orbital affects the size of the 2s orbital; the 2s must have its node fit precisely with the 1s so that there is zero overlap. If the 1s orbital contracted, the 2s orbital would also have to contract. In contrast, the 2p orbitals, because there is no “1p orbital,” have no such restriction and can lower their energy by contracting toward the nucleus.5,6

Fig. 1 Atomic orbitals of different angular momentum are always orthogonal with equal amounts of constructive and destructive interference. As an example, (top) a 1s orbital and 2p orbital on the same atom have equal amounts of constructive (unshaded with unshaded) and destructive (unshaded with shaded) interference. (Bottom) The 1s and 2p wavefunctions are plotted. Centered around the origin (nucleus), the constructive interference (right of the origin) is equivalent to the destructive interference (left of the origin).

4

Models for Understanding Main Group and Transition Metal Bonding

The result of the 2p-orbital contraction, and primogenic expansion of 2s, is that the 2s and 2p orbitals for the first-row p-block elements are essentially the same size. (For the rest of the p-block, the valence s-orbital is about 80% the size of p because both orbitals are primogenically expanded and p has a higher angular momentum.) Since the 2s and 2p orbitals are similar in size, they more effectively hybridize than, for example, 3s and 3p. Thus, the directional bonding and multiple bonding that dominate organic chemistry are due in large part to the primogenic effect. To illustrate, consider the very familiar geometry of acetylene versus Ar–GeGe–Ar (Fig. 2).7 While acetylene has the familiar linear geometry with sp hybridization on carbon, all the heavier congeners are bent. The experimental geometry of the germanium complex has Ge–Ge–Ar angles of 128.7(1) , closer to that expected for sp2 than sp. Natural Bond Orbital calculations on the H–GeGe–H model compound (Fig. 2) identify the hybridization of the orbital used by germanium to form the Ge–Ge s-bond as essentially sp; however, the Ge–H bond is formed by an sp2 hybrid, which gives a bent geometry. The primogenic effect limits the sp hybridization somewhat for the heavier congener, giving a preference for bonding with greater p character. (This is only part of the answer as the lower electronegativity of the elements below carbon also affects the hybridization, Bent’s Rule, vide infra.) There are many other common chemical observations that may be attributed to primogenic repulsion, or lack thereof. Of particular importance for our purposes here is the contraction of the 3d orbitals of the first row of the transition series. The 3d orbitals are smaller than might be expected from their heavier congeners. As a result, the 3d orbitals are only about 1/3 the size of the valence 4s-orbitals for these metals.8 In addition, the 3d-orbitals are very close to the same size as the core 3p orbitals; for example, the 3p-orbitals of iron have radiia of 0.373 A˚ , while the 3d orbitals have radii of 0.364 A˚ , slightly smaller than these core orbitals.8 So, as ligands approach the metal for bonding with the 3d orbitals, there is repulsion by core orbitals of essentially the same size. The relationship between the 3d and 3p orbitals for the first-row d-block elements is shown in Fig. 3. The early transition metals have d-orbitals slightly extended from the core, while the later metals have valence d-orbitals slightly smaller than these core orbitals. In contrast to the first-row transition metals, the second and third row elements have d orbitals that extend well beyond the core. For example, ruthenium has 4d orbital radii of 0.616 A˚ , while the 4p orbitals are smaller at 0.515 A˚ . The sizes of these orbitals are shown in Fig. 4, and the valence 4d-orbitals are always larger than the core orbitals for the elements. (The third-row metals are slightly larger but provide similar ratios: e.g., osmium has 5d-orbials with radii of 0.682 A˚ , with 5p radii of 0.550 A˚ .) An important consequence of primogenic repulsion on the valence orbitals for the 2nd and 3rd row transition elements, and the lack thereof for the 1st row, is that the bonds formed by the heavier congeners are stronger than the 1st row for similar compounds. The small size and core repulsion effects in the 1st row result in relatively weak bonds and weaker ligand fields than found for the 2nd and 3rd rows.

128.7(1)°

H

C

C

H

Carbon has sp-hybridization

Ar

Ge Ge

Ar

124.0°

Experimental geometry of Ge2Ar2

Ar

Ge Ar

Ge

Calculated geometry of Ge2H2

Ge

Ge

Fig. 2 (Top left) The familiar geometry of acetylene, which has sp-hybridized carbons due to good hybridization between 2s and 2p, a result of the primogenic effect. (Top center and bottom) Line drawing and ORTEP diagram from the X-ray diffraction data on Ar–GeGe–Ar, where Ar ¼ 2,6-(2,6-diisopropylphenyl)phenyl. (Top right) The calculated geometry of H–GeGe–H using M06L/aug-cc-PVTZ on G19. Ar ¼ 2,6-(2,6-Pri2C6H3)2C6H3. a Waber and Cromer calculated the “maximum in the charge density and its corresponding radius” for the atomic valence orbitals, which I’ll simply refer to as the “orbital radius.”8

Models for Understanding Main Group and Transition Metal Bonding

5

Fig. 3 The sizes of the valence 3d (red) and core 3p-orbitals (blue) for the first-row transition elements.8

Fig. 4 The sizes of the valence 4d (orange) and core 4p-orbitals (light blue) for the second-row transition elements.8

The primogenic effect dramatically affects the chemistry of the f-block as well. The f-orbitals of the lanthanides contract toward the nucleus so that they are less than 20% the size of the valence s-orbital, far smaller than some of the core orbitals. For example, for cerium, the f-orbitals are 0.366 A˚ , while the core 5p-orbital is over twice as large at 0.825 A˚ . As a result, the f-orbitals of the lanthanides are so small that they often cannot overlap with ligand orbitals and are rarely involved in bonding. In contrast, the actinides have primogenically expanded 5f-orbitals that may participate in bonding.8

1.02.3

The 2-center 2-electron heterocovalent bond and polar covalence theory

A series of concepts we are taught from our initial indoctrination as chemists regard covalent 2-center 2-electron (2c2e) bonds, such as those found in H2 and HCl. According to Molecular Orbital (MO) theory, championed by Mulliken, overlap leads to bonding and antibonding molecular orbitals where the lower energy bonding orbital is occupied. In Hybridization theory, championed by Pauling, the bond is viewed as a series of resonance forms, with some ionic contributions. In the end, the two viewpoints are mathematically equivalent at their limits,9 and the choice of one model over the other is based on expediency in addressing the problem at hand.

6

Models for Understanding Main Group and Transition Metal Bonding

Pauling was such an effective advocate that Valence Bond (VB) theory was regarded as the “correct” method for understanding chemical bonding until there was a series of instances where VB theory was thought to fail to give the correct answer where MO theory was successful. Common examples where VB theory was thought to fail are in describing the electronic structure of O2 and photoelectron spectrum of CH4. In actuality, these “failures” of VB theory are attributable more to poor use of the theory than the theory itself.10,11 A major advantage of VB theory is that the localized bonds are more easily correlated with the lines in our Lewis structures than delocalized molecular orbitals of MO theory. Despite this advantage, VB theory fell out of favor because MO theory was easier to adapt to computational methods. Computers were powerful allies to MO theory, leaving VB methods behind for many years. In some ways related to the orbitals used in VB theory, Natural Orbitals were first discussed by Löwdin in 1955, but Weinhold and Landis turned them into a powerful tool for modern chemists. Effectively, Natural Bond Orbital (NBO) theory allows the use of modern computation methods (DFT, coupled cluster theory, etc.) to adopt many advantages of VB theory. In this chapter, I will refer to “Hybridization theory” as a general term that incorporates the underlying concepts of VB and NBO theory. The underlying conceptual underpinnings of VB theory (and NBO theory) regain relevance as quantitative methods can be used in conjunction with the back-of-the-envelope calculations. In this section, it is advantageous to outline a Hybridization theory system for understanding heterocovalent bonds, Polar Covalence theory. The theory is not put forward as a replacement for accurate quantum mechanical calculations or experimental bond energies. Instead, it is a method for calculating bond energies by hand that illustrates several important concepts about bonding. No derivations of the equations are given to provide room for a discussion of concepts that underlie these descriptions of covalent bonds. Polar Covalence theory was developed by R. T. Sanderson as a method for the calculation of bond energies.12,b The method uses only a few descriptors for the elements and the bond distance, R0, in the compound to calculate the bond energy, typically within 5% of experimental values. The tabulated values are Sanderson electronegativity (wS), the change in electronegativity of an element with unit charge (DwS), covalent radii (rc), and the homonuclear bond energies of the elements. The homonuclear bond energies are described as being “fully unweakened” (E000 ), “partially weakened” (E00 ), and “fully weakened” (E0 ). A selection of these values is shown in Table 1. The cause of the weakening (difference between E000 , E00 , and E0 ) is a very interesting one that explains many chemical phenomena. Unfortunately, it will also require some building of foundations before the explanation can be given and the cause described. For the moment, we will call it Lone Pair Bond Weakening (LPBW), which was Sanderson’s name for the effect, an effect that wasn’t fully elucidated until 1996.15 On the surface, Sanderson electronegativity (wS) has nothing to do with bond energies to the element, unlike Pauling electronegativity, which is derived from bond energies. In contrast, Sanderson electronegativity is derived from electron density of the atoms (z) and their relationship to a calculated ideal electron density (zi). The electron density around the atom is simply as shown in Eq. (1). The electron density of an atom X is the atomic number divided by the volume of the sphere defined by the covalent radius of the atom in units of electrons/A˚ 3. Z e zX ¼
4 3 3 ˚ 3 pr A

(1)

The ideal electron density was found by making an interesting assumption. It was assumed that the noble gases were not only unreactive because of a closed shell, but also because they had an ideal electron density around the atom. The ideal electron density (zi) for non-noble gas atoms was found by linear interpolation between the nearest noble gases; in other words, the two noble gases that are the bookends for an atomic shell were used to define a line of electron density relative to atomic number that was ideal for the atoms in between. Sanderson electronegativity is simply the ratio between the electron density of an atom (z) and the ideal electron density (zi), i.e., wS ¼ (z)/(zi). Originally, these values were called “stability ratios” as they gave a measure of how the atom could be stabilized by addition or subtraction of electron density, and fluorine had a value of 5.75 on this scale. Later, Sanderson scaled all his values to the more familiar Pauling value of wF ¼ 4.00, and the numbers became known as Sanderson electronegativity.16 At this stage, the equations for determination of bond energies using Polar Covalence theory will be given (Table 2) and a brief explanation of each will be provided with the underlying concept. As mentioned, Polar Covalence theory is based on Hybridization theory and suggests that the bond in a heterocovalent compound can be described using only two resonance forms (Fig. 5). In the left form, there is a purely covalent bond between atoms A and X, signified by a line. In the right form, there is a purely ionic form where the less electronegative atom (A) has a positive charge, and the more electronegative atom (X) has a negative charge. An alternative, minor, resonance form where the less electronegative atom has a negative charge is disregarded. The overall bond energy (the average enthalpy of the bonds in the compound) according to Polar Covalence theory is comprised of the covalent bond energy, Ec, and the ionic bond energy, Ei, with two weighting coefficients tc and ti. The sum of the contributions from all resonance forms should be unity, and these are the only forms being considered, so tc +ti ¼ 1 (Eq. 3 in Table 2). The overall b

Here, “bond energy” refers to the average bond enthalpy for homolytic cleavage of the bonds in a compound, e.g., 1/3 of the enthalpy required to atomize BF3 to B + 3 F. A more detailed explanation and examples are given below.

Models for Understanding Main Group and Transition Metal Bonding

7

Table 1 Sanderson electronegativities (wS), change in electronegativity with unit charge (D wS), covalent radii (rc) in A˚ ( 10−10 m), homonuclear fully unweakened (E000 ), partially weakened (E00 ), and fully weakened (E0 ) bond energies in kcal/mol.a Element

wS

H Li Be B C N O F Na Mg Al Si P S Cl K Ca Sc(II) Sc(III) Ti(II) Ti(III) Ti(IV) V(II) V(III) V(IV) V(V) Cr(II) Cr(III) Cr(IV) Cr(V) Cr(VI) Mn(II) Mn(III) Mn(IV) Mn(V) Mn(VI) Fe(II) Fe(III) Co(II) Co(III) Co(IV) Ni(II) Ni(III) Ni(IV) Ni(V) Cu(II) Zn Ga Ge As Se Br Rb Sr Y(II) Y(III) Zr(II) Zr(III) Zr(IV) Nb(II) Nb(III) Nb(IV)

2.592 0.670 1.810 2.275 2.746 3.194 3.654 4.000 0.560 1.318 1.714 2.138 2.515 2.957 3.475 0.445 0.946 0.64 1.02 0.73 1.09 1.50 0.69 1.39 1.89 2.51 1.24 1.66 2.29 2.83 3.37 1.66 2.20 2.74 3.28 3.82 1.64 2.20 1.96 2.56 3.10 1.94 2.73 3.27 3.81 1.98 2.223 2.419 2.618 2.816 3.014 3.219 0.312 0.721 0.40 0.65 0.52 0.79 0.90 0.77 1.02 1.25

 pffiffiffiffiffi DwS ¼ 1:57 wS

2.528 1.285 2.112 2.368 2.602 2.806 3.001 3.140 1.275 1.802 2.055 2.296 2.490 2.700 2.927 1.047 1.527 1.256 1.586 1.341 1.639 1.923 1.304 1.851 2.158 2.487 1.748 2.023 2.376 2.641 2.882 2.023 2.329 2.599 2.843 3.069 2.011 2.329 2.198 2.512 2.764 2.187 2.594 2.839 3.065 2.209 2.341 2.442 2.540 2.635 2.726 2.817 0.877 1.333 0.993 1.266 1.132 1.395 1.489 1.378 1.586 1.755

Rc

E000

0.320 1.336 0.887 0.822 0.772 0.734 0.702 0.681 1.539 1.373 1.258 1.169 1.107 1.049 0.994 1.962 1.74 1.695 1.695 1.650 1.650 1.650 1.605 1.605 1.605 1.605 1.560 1.560 1.560 1.560 1.560 1.516 1.516 1.516 1.516 1.516 1.471 1.471 1.426 1.426 1.426 1.381 1.381 1.381 1.381 1.336 1.292 1.256 1.223 1.294 1.167 1.142 2.16 1.91 1.868 1.868 1.826 1.826 1.826 1.784 1.784 1.784

104.2 24.6 67.6 76.7 85.4 94.9 104.0 113.1 16.4 42.3 48.2 54.1 60.0 65.9 71.8 13.1 30.8 31.3 31.3 31.8 31.8 31.8 32.3 32.3 32.3 32.3 32.8 32.8 32.8 32.8 32.8 33.3 33.3 33.3 33.3 33.3 33.8 33.8 34.3 34.3 34.3 34.8 34.8 34.8 34.8 35.3 35.8 40.3 44.8 49.3 53.8 58.3 12.4 24.6 25.1 25.1 25.7 25.7 25.7 26.3 26.3 26.3

E00

E0

66.9 68.8 76.8

38.8 33.6 38.2b

56.9 60.4 64.9

53.7 54.9 58.0

45.1 46.0 52.2

40.9 38.2 46.1

(Continued )

8

Models for Understanding Main Group and Transition Metal Bonding

Table 1

(Continued)

Element

wS

Nb(V) Mo(II) Mo(III) Mo(IV) Mo(V) Mo(VI) Cd In Sn(IV) Sn(II) Sb Te I Cs Ba Hf(II) Hf(III) Hf(IV) Ta(II) Ta(III) Ta(IV) Ta(V) W(II) W(III) W(IV) W(V) W(VI) Tl(III) Tl(I) Pb(IV) Pb(II) Bi

1.42 0.90 1.15 1.40 1.73 2.20 1.978 2.138 2.298 1.477 2.458 2.618 2.778 0.220 0.651 0.31 0.56 0.81 0.44 0.69 0.94 1.17 0.73 0.98 1.23 1.48 1.67 2.246 0.98 2.291 1.900 2.342

 pffiffiffiffiffi DwS ¼ 1:57 wS

1.871 1.489 1.684 1.858 2.065 2.329 2.208 2.296 2.380 1.908 2.461 2.540 2.617 0.736 1.267 0.874 1.175 1.413 1.041 1.304 1.522 1.698 1.341 1.554 1.741 1.910 2.029 2.353 1.560 2.376 2.164 2.403

Rc

E000

1.784 1.742 1.742 1.742 1.742 1.742 1.493 1.455 1.420 1.420 1.389 1.360 1.333 2.35 1.98 1.80 1.80 1.80 1.76 1.76 1.76 1.76 1.73 1.73 1.73 1.73 1.73 1.490 1.49 1.48 1.48 1.47

26.3 26.9 26.9 26.9 26.9 26.9 30.4 33.1 35.8 35.8 38.5 41.2 43.9 10.8 22.2 15.6 15.6 15.6 14.7 14.7 14.7 14.7 13.8 13.8 13.8 13.8 13.8 16.6 16.6 24.2 24.2 32.2

E00

E0

35.8 39.2 40.0

33.0 37.2 36.1

a

Values are from Sanderson’s references on Polar Covalence.12,13 Value from Huber and Herzberg.14

b

Table 2

Polar Covalence Theory equations for determining homolytic bond energies.

Determination of the A–X bond energy Sum of the ionic and covalent coefficients is unity Covalent bond energy

EA−X ¼ tcEc + ttEt 1 ¼ tc + tt pffiffiffiffiffiffiffiffiffiffiffiffiffi E c ¼ b RR0c E AA E XX Rc ¼ RA + RX Bond order Single Double Triple

Ionic bond energy

(2)a (3) (4)b

b 1.000 1.488 1.787

˚

=mol E i ¼ b  332 kcalA ˚

(5)

wM −wA Þ ðw −w Þ pffiffiffiffi ¼ M A dA ¼ ð1:57 wA DwA qffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi ffi A B Z N N N wm ¼ wA  wB  ⋯  wNZ

(7)d

R0 A

Estimation of the ionic coefficient Charge on an atom in a molecule Molecular electronegativity a

t t ¼ jdA j 2+jdX j

(6)c (8)e

tc ¼ covalent coefficient, ti ¼ ionic coefficient, Ec ¼ covalent bond energy, Ei ¼ ionic bond energy. Rc ¼ covalent bond distance from sum of covalent radii (RA + RX), R0 ¼ bond distance, EAA ¼ covalent bond energy of A, EXX ¼ covalent bond energy of X. c d ¼ partial charge. d wM ¼ molecular electronegativity, wA ¼ electronegativity of A. e N ¼ number of total atoms, NA ¼ number of atoms A. b

Models for Understanding Main Group and Transition Metal Bonding

9

Fig. 5 Resonance forms considered in Polar Covalence theory.

bond energy in A–X is then EAX ¼tcEc +tiEi (Eq. 2 in Table 2). Since tc and ti are related by Eq. (3) in Table 2, by finding only three values—Ec, Ei, and ti—we can calculate the bond energy for the system. All three values are relatively simple to calculate, as will be shown. First, Sanderson showed that the apparent bond order (the number of bonds a chemist would normally draw for the compound, e.g., a single bond for H2 and a double bond for O2) has the effect of adding a scalar (b) to Ec and Ei. As any chemist will tell you, a double bond is not twice as strong as a single bond, and a triple bond is not three times stronger than a single bond. Sanderson found for a double bond that b ¼ 1.488 and for a triple bond b ¼ 1.787. There is, however, an additional correction for the bond shortening that occurs in these higher bond orders that leads to some strengthening over these factors (vide infra). Sanderson proposed that the covalent bond energy (Ec) was simply the geometric mean of the homolytic bond energies for the elements involved. For example, the purely covalent bond energy of NaCl would be the geometric mean of the Na2 and Cl2 bond energies. In essence, this suggests that there is a primal energy associated with a specific atom involved in purely covalent bonding that is fully manifested when the atom bonds to itself. When the atom bonds to an atom of a different type, e.g., A–X, then the covalent bond energy is the geometric mean of these primal covalent energies of A and X. (This same assumption was used by Pauling in the generation of his electronegativity values, without the corrective value for the bond distance. There are some caveats to this having to do with LPBW effect, which will be discussed in due course.) To this there was applied a correction involving the ratio of the actual bond distance in the compound and the sum of the covalent radii. Shorter bond distances increase overlap and generally increase bond energy. This, in addition to the factor for the bond order (b), is what leads to the equation for Ec (Eq. 4 in Table 2). The ionic bond energy Ei is very simple and comes directly from Coulomb’s Law. The ionic bond energy (Eq. 5 in Table 2) is simply a constant (332 kcal A˚ /mol) divided by the bond distance, R0. This equation gives Ei in kcal/mol if the bond distance is in A˚ . So far, two of the three values (Ec, Ei, and ti) needed to calculate the bond energy have been found; all that remains is the ionic weighting coefficient, ti. The contribution of the ionic form to the overall electronic structure is related to the partial charges in the system. If there are larger partial charges on A and X, then ti should be larger. If there is a full positive charge on A (dA ¼ +1) and a full negative charge on X (dX ¼ −1), then ti ¼ 1 and only the ionic form contributes. As a result, ti can be estimated as the average of the absolute values of the partial charges on the two atoms, i.e., (|dA | + | dX |)/2 ¼ ti (Eq. 6 in Table 2). To find ti, we need the partial charges on the atoms, which are related to the electronegativities of the atoms involved. There are many charge schemes that have been developed over the years using computational methods. Here, Sanderson developed a simple method for calculating charges in a molecule by hand. The method relies on a basic principle of covalent systems, electronegativity equalization.12 Consider a simple diatomic in the gas phase, e.g., HCl. Chlorine is more electronegative than hydrogen, meaning it will attract more electron density in the bond. As a result, the chlorine atom accumulates some negative charge, and the hydrogen accumulates some positive charge. A partial positive charge makes hydrogen more electronegative than neutral hydrogen, and a partial negative charge makes chlorine less electronegative than neutral chlorine. In other words, charge transfers between the two atoms until the electronegativities of H and Cl in HCl equalize. Electronegativity equalization simply suggests that in a symmetric covalent system, the electronegativity of all the atoms is the same. Sanderson’s simple charge scheme relies on electronegativity equalization to find the molecular electronegativity, and the charges on individual atoms are then related to the difference between the molecular electronegativity and the electronegativity of the individual atoms. The electronegativity of the molecule (wM) is found by simply taking the geometric mean of the electronegativities of all the atoms in the molecule (Eq. 8 in Table 2). The partial charge on an individual atom (dA) is then the difference between the molecular electronegativity (wM) and the atomic electronegativity (wA) divided by the change in electronegativity per unit charge (DwA) for that element. Sanderson found empirically that the change in electronegativity with unit charge was equal to 1.57 times the square root of the neutral atom’s pffiffiffiffiffi electronegativity, i.e., DwA ¼ 1.57 wA . The electronegativity would go up with a positive charge by that amount and down with a negative charge. Using fluorine as an example, wF ¼ 4.000, and DwF ¼ 1.57(2) ¼ 3.140. In other words, the electronegativity of F+ would be 7.140, but F− has an estimated electronegativity of 0.86. It is important to note the limitations of the charge scheme. First, this simple scheme does not work for molecules that have the same atom in two very different environments, e.g., H3Si–SiF3. The scheme is too primitive to distinguish between the different charges on silicon in H3Si– and –SiF3. Second, this charge scheme does not work for molecules with dative bonds. (Much more on this later.) As a rule of thumb, this scheme works best for binary, covalent compounds with an AXn formula. Despite its limitations, there are thousands of compounds where this simple method can be applied very successfully, and it illustrates many bonding principles.

10

Models for Understanding Main Group and Transition Metal Bonding

As an example of using the charge scheme, consider the molecule TiCl4. To find the charges, one needs to find the molecular electronegativity based on the atoms in the system (wTi(IV) ¼ 1.50, wCl ¼ 3.475). The molecular electronegativity is the geometric mean of the electronegativity of the atoms due to electronegativity equalization. qffiffiffiffiffiffiffiffiffiffiffiffi qffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi 5 wTiCl4 ¼ 5 w4Cl wTi ¼ ð3:475Þ4 ð1:50Þ ¼ 2:938

The charge on any individual atom is then found using Eq. (7) in Table 2, where the difference in electronegativity between the molecule and atom is divided by the change in electronegativity per unit charge. dCl ¼ dTi ¼

wTiCl4 − wCl 2:938 − 3:475 ¼ − 0:183 ¼ 2:927 DwCl wTiCl4 − wTi 2:938 − 1:50 ¼ + 0:748 ¼ DwTi 1:923

The contribution of the ionic resonance form is the average of the absolute values of the charges across one of these bonds (Eq. 6 in Table 2). ti ¼

jdCl j + jdTi j 0:183 + 0:748 ¼ ¼ 0:466 2 2

In other words, the bonding in TiCl4 is about 47% ionic. The remainder, of course, is covalent, and tc ¼ 0.534. For TiCl4, we have calculated the weighting coefficients ti and tc. To calculate the bond energy for the compound, we simply need to calculate Ei and Ec, the ionic and covalent bond energies, respectively. The only piece of data outside of Table 1 necessary to do this is the Ti–Cl bond distance, R0, which is 2.185  0.010 A˚ from X-ray diffraction.17 Using this, we can readily calculate both Ei and Ec using Eqs. (4) and (5) in Table 2. For the covalent bond energy, we will use the “fully unweakened” value for chlorine; the reason for this will be explained shortly. Notice that the ionic bond energy is much larger than the covalent bond energy, and this is almost always the case. In other words, ionic bonds are typically stronger when comparing related systems, a point that will be illustrated further in a moment. qffiffiffiffiffiffiffiffiffiffiffiffi qffiffiffiffiffiffiffiffiffiffiffiffi ðrTi + rCl Þ 000 000 000 000 ETi ECl ETi ECl ¼ R0 rffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi 1:650 A˚ + 0:994 A˚ kcal kcal ¼  71:8 ¼ 57:8 kcal=mol 31:8 ˚ mol mol 2:185 A Ec ¼ b

Rc R0

˚

Ei ¼ b 

˚

A A 332 kcal  mol 332 kcal  mol ¼ ¼ 151:9 kcal=mol R0 A˚ 2:185 A˚

The mean bond energy of TiCl4 is then found using Eq. (2) in Table 2 as being 101.7 kcal/mol. This can be compared with the experimental mean bond energy of TiCl4, which is 102.7 kcal/mol, an error of 1%. Keep in mind that the calculation only required the values in Table 1 and the experimental bond distance to get this accuracy and, with a little practice, can be done in a few minutes with just a calculator.     kcal kcal + 0:534 57:8 ¼ 70:8 + 30:9 ¼ 101:7 kcal=mol ETiCl4 ¼ ti Ei + tc Ec ¼ 0:466 151:9 mol mol One quick point that needs to be made is the difference between bond energy and bond dissociation enthalpy (BDE). The IUPAC definition of bond energy is, “The energy required to break a given type of bond between atoms in certain valence states. An averaged bond energy is commonly derived by dissecting the heat of atomization of a molecule into contributions of individual bonds. For molecules with localized bonds, the heats of atomization (formation) are usually well approximated by the sum of pertinent averaged bond energies.” In other words, the bond energy is the average enthalpy required to break the bonds in a molecule. The “bond energy” or, better, “mean bond energy,” of methane is ¼ the energy required to atomize CH4(g) to C(g) + 4 H(g). Similarly, the bond energy in TiCl4 is one-fourth the enthalpy of the reaction TiCl4(g) !Ti(g) + 4 Cl(g).

Models for Understanding Main Group and Transition Metal Bonding

11

This should not be confused with BDE, which is (IUPAC), “the enthalpy (per mole) required to break a given bond of some specific molecular entity by homolysis, e.g., CH4 !CH3 + H”. In other words, BDE is the enthalpy to break a specific bond while (mean) bond energy often refers to the average enthalpy required to break all the bonds of a type in a compound. To provide a further illustration, consider the bond energy in TiCl3. (Would you expect the bond energy in this Ti(III) complex to be higher or lower than in the Ti(IV) complex just discussed?) Using Polar Covalence theory, we can quickly calculate the bond energy in TiCl3 using the data in Table 1 and using a Ti–Cl bond distance of 2.197 A˚ from DFT calculations at the M06L/cc-PVTZ level. Calculation of Ec and Ei goes as shown below. Notice that both Ec and Ei are very similar to the Ti(IV) case above, but both are slightly smaller due to the increased bond distance. qffiffiffiffiffiffiffiffiffiffiffiffi qffiffiffiffiffiffiffiffiffiffiffiffi ðrTi + rCl Þ 000 000 000 000 ETi ECl ETi ECl ¼ R0 rffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi 1:650 A˚ + 0:994 A˚ kcal kcal  71:8 ¼ ¼ 57:5 kcal=mol 31:8 mol mol 2:197 A˚ Ec ¼ b

Rc R0

˚

Ei ¼ b 

˚

A A 332 kcal  mol 332 kcal  mol ¼ ¼ 151:1 kcal=mol R0 A˚ 2:197 A˚

The charges of the atoms and ti are calculated as below. (The charges don’t quite cancel, but the difference is a few thousandths when dividing the central atom charge by 3.) qffiffiffiffiffiffiffiffiffiffiffiffi pffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi 4 wTiCl3 ¼ 4 wTi w3Cl ¼ 1:09  3:4753 ¼ 2:601 dCl ¼ dTiðIIIÞ ¼

ti ¼

wTiCl3 − wCl 2:601 − 3:475 ¼ − 0:299 ¼ DwCl 2:927

wTiCl3 − wTiðIIIÞ 2:601 − 1:09 ¼ ¼ + 0:922 DwTiðIIIÞ 1:639

jdCl j + dTiðIIIÞ 0:299 + 0:922 ¼ 0:611 ¼ 2 2 t c ¼ 1−0:611 ¼ 0:389

Importantly, the ionic and covalent energies are slightly smaller here for Ti(III) as mentioned, but titanium(III) has a lower electronegativity, which leads to a much more ionic compound. The titanium(III) complex is 61% ionic versus TiCl4 being 47% ionic. The calculation of the bond energy of TiCl3 is shown below.     kcal kcal + 0:389 57:5 ¼ 92:3 + 22:4 ¼ 114:7 kcal=mol ETiCl3 ¼ ti Ei + tc Ec ¼ 0:611 151:1 mol mol Even though the Ti–Cl bond length in titanium(III) chloride is longer than in TiCl4, the bond energy is higher because the compound is more ionic. As mentioned, more ionic bonds are typically stronger, and the DFT calculated bond energy for TiCl3 (M06L/cc-PVTZ) is 119.3 kcal/mol, a difference of 4% from Polar Covalence Theory. One of the most important ligands for organometallic chemistry is carbon monoxide, which is a ligand we will examine later with regards to the Lone Pair Bond Weakening effect. In addition, calculation of the bond energy in CO provides an example with a multiple bond. The experimental bond distance (R0) in CO is 1.128 A˚ . In this instance we will use the “fully unweakened” value (R000 ) for the O–O bond to calculate the covalent bond energy. qffiffiffiffiffiffiffiffiffiffiffi Rc 1:474 pffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi 000 000  85:4  104 ¼ 1:787  1:307  94:2 ¼ 220:1 kcal=mol EC EO ¼ b  Ec ¼ b 1:128 R0 Ei ¼ b 

332 kcal  A˚ =mol ¼ 1:787294:3 kcal=mol ¼ 525:9 kcal=mol 1:128 A˚

It is obvious from Ec and Ei that the bond energy is quite large. Now, we need a value for ti. The molecular electronegativity (wCO) is simply the geometric mean of wO and wC, 3.168. From this, and the values in Table 1, we can calculate the charges on carbon and oxygen. The average of the absolute values of two charges is then ti ¼ 0.162, and tc ¼ 0.838.

12

Models for Understanding Main Group and Transition Metal Bonding

wCO ¼

pffiffiffiffiffiffiffiffiffiffiffi pffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi wC wO ¼ 2:746  3:654 ¼ 3:168

dC ¼

wCO − wC 3:168 − 2:746 ¼ ¼ + 0:162 DwC 2:602

dO ¼

wCO − wO 3:168 − 3:654 ¼ − 0:162 ¼ 3:001 DwO

ti ¼

jdC j + jdO j j +0:162j + j −0:162j ¼ ¼ 0:162 2 2

The bond energy is then the sum of the weighted values for the ionic and covalent bond energies, which gives a value of 269.6 kcal/mol. The experimental bond energy of CO is 256 kcal/mol, which means Polar Covalence has an error of 5% in this case. ECO ¼ tc Ec + ti Ei ¼ 0:838  220:1 + 0:162  525:9 ¼ 269:6 kcal=mol Simple Polar Covalence theory illustrates many different principles that helps to build intuition regarding bonding: electronegativity equalization, bond energy variance with polarity, effect of bond distance on bond energy, covalent vs ionic bonding, nature of partial charges relative to molecular electronegativity, and others. Even in cases where Polar Covalence calculations are not applicable, these principles still can steer the chemist toward the correct conclusions. For example, Mills and coworkers published the reaction below (Eq. 9).18 Polar Covalence calculations cannot be applied directly to Me2P–P(CF3)2; the charge scheme fails because phosphorus appears in two very different environments. However, the principle we illustrated above—that more polar bonds are more stable for similar systems—suggests that the reaction should proceed to completion, which is experimentally the case. The mixed CH3/CF3 system should have the highest P–P bond ionic character and be favored. Computationally, it is found that DHrxn ¼ −9 kcal/mol (B3LYP/aug-cc-pvdz). ð9Þ

1.02.4

The 2-center 2-electron dative bond

The other, just as important, type of commonly encountered 2c2e bond is the dative bond, a Lewis acid-base interaction. Two examples of these bond types are found in Fig. 6, where CuCl is used as the archetypal heterocovalent bond and H3NBH3 for a dative bond. The heterocovalent bond is signified by a simple line between the two atoms, and the most important resonance forms are the ones shown: one where there is a purely covalent bond between Cu and Cl with equal electron sharing and another where the bond is polarized so that the more electronegative chlorine has a negative charge. (A third resonance form containing Cu− and Cl+ would be expected to make a very small contribution and is ignored for the sake of discussion.) Often, chemists will differentiate between the covalent and dative bonds in their drawings. The dative bonds are often signified with an arrow from the Lewis basic atom to the acidic atom as shown on the top right in Fig. 6. Occasionally, formal charges are added as on the bottom right in the figure. For expediency, inorganic and organometallic chemists often use a simple line for dative bonds as well, Fig. 7, even though the lack of distinction from covalent bonds may lead to confusion. (Particularly troubling is the use of a line for the bond between an N-heterocyclic carbene and a metal center without use of the formal charges and the double bond between nitrogen and carbon, which leads the uninitiated to assume there is a hydrogen on the bonding 3-coordinate carbon when there is none.) If we draw dative bonds differently and think of them differently from heterocovalent bonds, in what ways do these two distinctive 2-center 2-electron bonds differ in a Hybridization theory picture? In looking at the electron density in a very polar heterocovalent bond and a dative bond, there may be little difference. In fact, Weinhold and Landis state, “Of course, there is intrinsically no sharp physical distinction between a strongly delocalized lone pair and a highly polarized dative bond.”19 (For context, this sentence appears in a section entitled, “Ionic-covalent transitions” focused on electron density in bonds.) One can see this in the s-bonds for CuCl and H3NBH3. On the left in Fig. 8 is the calculated s-bonding NBO for CuCl, which is strongly polarized toward the chloride as expected from the large electronegativity difference between the two atoms. If one considers what the structure would look like if the bonding pair

Fig. 6 (Left) CuCl as an example of a heterocovalent bond that can be thought of as being comprised of covalent and ionic resonance forms. (Right) H3NBH3 as an example of a dative bond where a major contributor is the shared electron form shown with charges on either side of the bond.

Models for Understanding Main Group and Transition Metal Bonding

13

Fig. 7 (Top) In some complexes with many dative bonds, the arrows are replaced by lines, dashes, and wedges, which better illustrates the stereochemistry. However, it is understood in these cases that the bonds are dative. (Bottom) N-Heterocyclic carbenes (NHCs) are examples in which the use of simple lines leads to particular confusion. If a line is drawn to the metal with no charges or lone pairs shown, there can be misunderstandings as to whether there is a hydrogen on the donor carbon. Two clearer methods for drawing NHC metal complexes are shown on the right.

H Cu

H H N

Cl

H

Cl–

Cl Cu

B

H H

BH3

BH3

Cu+

H3N

H3N



Fig. 8 The s-bonding NBOs for CuCl and H3NBH3 (NBO7, B3LYP/6-311++G ). At the top are line drawings with the same orientation as the calculated structures in the middle for reference. At the bottom are resonance forms implied by the highly polarized 2-center 2-electron bonds.

of electrons were localized on the more electronegative atom, one gets a resonance form of Cu+ Cl−, the expected ionic form. A large contribution by the form where we polarize the electrons to one side strengthens the bond through increased inclusion of the polar form (vide supra). On the right is the calculated s-bonding NBO for H3NBH3. Again, the bond is highly polarized toward one atom, in this case toward nitrogen. However, if we move toward the extreme where we localize more of the bonding pair of electrons on the more electronegative nitrogen atom, we get less of a bond between N and B. In other words, a large contribution by the form where we polarize the electrons to one side weakens the bond in the dative case and strengthens the bond in the covalent case. Defining the precise difference between a dative and covalent 2-center 2-electron bond is a difficult task. As is almost always the case in chemistry, it is difficult to make an assertion that has no exceptions. Chemistry is too broad a field, and chemists are too clever, for there not to be exceptions found for any rule put forward. Nevertheless, a few definitions of a dative 2c2e bond will be discussed as they do relate to topics already put forward and are somewhat enlightening to the properties of the bonds. One of the more common definitions for 2-center 2-electron bonds is an operational one: dative bonds tend to heterolytically cleave to reform the Lewis acid and base, whereas covalent bonds tend to cleave homolytically to form radical species.c In addition, generally speaking, dative bonds are weaker than similar heterocovalent bonds. If we consider the electronic structure of heterocovalent bonds to be largely comprised of covalent and ionic resonance contributors and dative bonds to be comprised of covalent c An exception that is often given to this covalent/dative dichotomy is the comparison of F2C]CF2 and H2C]CH2. Cleaving F2C]CF2 gives singlet-F2C, while H2C]CH2 gives triplet-H2C. Since both are covalent by the usual reckoning, it is thought that because one gets a lone pair on the carbon after cleaving the bond there must be dative character. However, the bond in question isn’t a 2c2e bond; the C–C bonds in ethylene and tetrafluoroethylene are 2-center 4-electron bonds, which makes them quite different species subject to spin state differences on cleavage. The theoretical work done on them is fascinating, but the question of them being dative or covalent is complicated by the fact that the C–C bonds in question are not 2c2e bonds at all.

14

Models for Understanding Main Group and Transition Metal Bonding

and unbonded contributors (Fig. 8), it is then unsurprising that dative bonds are typically weaker than similar covalent bonds. Consider the two isoelectronic species H3NBH3 and H3CCH3. The covalent C–C bond in ethane has a BDE of 89 kcal/mol, while the heterolytic bond energy of H3NBH3 is much weaker at 31 kcal/mol. If a different reference than ethane is preferred, the covalent bond in H2BNH2 has a homolytic bond energy of 141 kcal/mol (G4 method), where the polarity in this bond increases the covalent bond energy over ethane. Another method of differentiating a dative bond from a covalent bond is dependent upon the electron-density distribution. In Polar Covalence theory, it was mentioned that the theory was not applicable to dative interactions; the charge scheme, which is dependent upon electronegativity equalization, fails for dative bonds. Creation of the dative bond from the Lewis acid/base pair leads to formal charges across the bond in the resonance form where the electrons are equally shared (Fig. 8). One can even see this charge difference using Natural Population Analysis. The Natural charges on H3BNH3 were calculated using NBO from electron-density determined using DFT at the B3LYP/6-311 + +G level. The neutral, isolated molecules BH3 and NH3 accumulate net charges consistent with expectations from Lewis structures, i.e., BH3 is negatively charged and the NH3 is positively charged (− 0.35 on BH3 and +0.35 on NH3, Fig. 9). Note that the charges run counter to the relative electronegativity of the fragments. (The Polar Covalence molecular electronegativity of BH3 is 2.509, and for NH3 the molecular electronegativity is 2.731.) There is something decidedly different about the bonding in this dative system and heterocovalent 2c2e bonds in the electron density distribution. Weinhold and Landis have explained this phenomenon by defining “Natural Electronegativity.” The ionicity of a bond in NBO is defined in Eq. (10), where cA and cX are the NBO polarization coefficients for atoms A and X. The ionic character in a bond (bionic AX ) is cov cov then bionic AX ¼ bAXiAX, where bAX is the bond order. The covalent contribution (bAX) is then bAX ¼ bAX(1 − iAX). iAX ¼

c2A − c2X c2A + c2X

(10)

They defined their electronegativity using Pauling-inspired Eq. (11). The ionic character in the bond is related to the electronegativities of the atoms involved. This equation was derived so that it placed atoms on the familiar 4.0 scale of Pauling and held well for covalent bonds. However, for dative bonds, it was found that the ionicity was not related to these NBO-derived electronegativities (Eq. 12), in synergy with the Polar Covalence theory arguments discussed earlier.


NBO  icov − wNBO (11) AB ’ 1 − exp − 0:45 wA X


 NBO − wNBO (12) idat AB ≄1 − exp − 0:45 wA X

In summary, both 2-center 2-electron dative and polar covalent bonds can have electron densities that are strongly polarized to one side. Nevertheless, the two bond types differ significantly. In a heterocovalent bond, the charges are largely set by electronegativity of the atoms involved. In a dative bond, the Lewis base has a higher (more positive) charge than expected based on electronegativity and the Lewis acid has a lower (more negative) charge than expected based on electronegativity considerations alone. These effects can be understood as consequences of the simple resonance forms for 2c2e bonds shown in Fig. 8.

1.02.5

Complementarity of qualitative hybridization theory and molecular orbital theory

Both Valence Bond and Natural Bond Orbital theories have a basis in “Hybridization theory.” While Molecular Orbital theory is more familiar to most organometallic chemists, this section will show that Hybridization theory can be quite complementary. As in previous sections, the focus will be on simple calculations that can be done in a few minutes with a calculator, but NBO calculations can be used to supplement the more intuitive analysis here. Some comparisons will be made between MO and Hybridization arguments as well. The focus in this section will be entirely on the main group, and a later section will address Hybridization theory of the transition metals. For any element in the main group, the valence orbitals are comprised of an s and three p orbitals, which can form sp hybrids. A 1st-row p-block hybrid (h) is formed as shown in Eq. (13).

Fig. 9 Natural charges in H3BNH3 at the equilibrium geometry (NBO7, B3LYP/6-311++G ).

Models for Understanding Main Group and Transition Metal Bonding 1 hsp ¼ pffiffiffiffiffiffiffiffiffiffiffiffiffi ½2s + að2pÞ 1 + a2

15

(13)

The electron density is related to the square of the wavefunction (Born’s rule), so we can square the function in Eq. (13) and integrate over all space (dt). The term “2a(2s)(2p)” involves the product of two different atomic orbitals, 2s and 2p, and the overlap integral must be zero because all the atomic orbitals are orthogonal by necessity (vide supra).d The middle integral is then dropped, and the probability function for the hybrid is redefined, Eq. (14), where the substitution of a2 ¼ l was also made. Z Z  1

h2sp dt ¼ ð2sÞ2 + 2að2sÞð2pÞ + a2 ð2pÞ2 dt 1 + a2 Z

Z Z Z 1 ð2sÞ2 dt + 2að2sÞð2pÞdt + a2 ð2pÞ2 dt h2sp dt ¼ 2 1+a Z Z Z 1 a2 2 h2sp dt ¼ ð 2s Þ ð2pÞ2 dt dt + 1 + a2 1 + a2 Z Z Z 1 l h2sp dt ¼ ð2sÞ2 dt + ð2pÞ2 dt (14) 1+l 1+l The term l is the “hybridization parameter,” which is the familiar exponent in spl, e.g., sp3 for l ¼ 3. When defined in this way, the parameter has no requirement to have integer values as we typically learn in organic classes. The hybridization parameter, l, is the ratio of p- to s-character in the hybrid. p − character l=ð1 + lÞ ¼ ¼l s − character 1=ð1 + lÞ The fraction of s-character in the hybrid is then related to 1/(1 + l). In fact, if one takes all the hybrids on an atom and adds up the 1/ (1 + l) factors for each hybrid, it must equal the number of valence s-orbitals, i.e., unity. This is called the “Sum Rule” for hybrids, which simply states that the hybrids on a main group atom sum up to the valence orbitals (one s-orbital and three p-orbitals), no more or less. Mathematically, they could be expressed as in Eq. (15) for the s-orbital and Eq. (16) for the p-orbitals, where the sum is over all the i hybrids on the atom. For our purposes here, we only need one Sum Rule or the other, and we’ll use the one for the s-orbital below (Eq. 15). X 1 ¼ 1 ðSum Rule for sÞ 1 +li i X li ¼ 3 ðSum Rule for pÞ 1 +li i

(15)

(16)

Another important point regarding hybrids is that, to maintain orthogonality, hybrids of a specific type must be at a specific angle to one another. In other words, two sp2-hybrids are always 120 apart. What if the hybrids are different? For example, what is the angle between sp- and sp3-hybrids on the same atom? The angle between any two hybrids required to maintain orthogonality is called the Coulson Directionality Theorem, which is shown in Eq. (17), where wij ¼ the angle between the i-th and j-th hybrid. 1 cos oij ¼ − pffiffiffiffiffiffiffiffi ðCoulson Directionality TheoremÞ li l j

(17)

For our hypothetical question above, we calculate that the angle between an sp1 and sp3 hybrid must be 120 : 1 1 1 cos oij ¼ − pffiffiffiffiffiffiffiffi ¼ − pffiffiffiffiffiffiffiffiffi ¼ − 2 li l j 13   1 ¼ 120 oij ¼ arccos − 2

If the two hybrids in question have the same hybridization parameter, then the Coulson Directionality Theorem simplifies to Eq. (18). For two sp3 hybrids, for example, one then gets arccos (−1/3) ¼ 109.47 , the tetrahedral angle. Looking at the limits of the hybridization parameter of 0 and infinity gives that sp0 is a pure s-orbital, and that sp1 is a pure p-orbital. cos oi¼j ¼ −

1 l

(18)

One can also define the energy of a main group hybrid (Eh) as a function of the s- to p-orbital energy gap (DEsp), as shown in Eq. (19).

d

We’re also assuming that all the functions are real.

16

Models for Understanding Main Group and Transition Metal Bonding

Eh ¼

l DEs!p 1+l

(19)

An important concept in Hybridization theory regards the nature of hybrids used to bond to different groups. In organic chemistry, hybridization is usually described as a kind of average hybridization for the carbon, e.g., the carbon in ethylene is described as “sp2.” In actuality, an atom will use a different hybrid for every different group on the atom. For example, in ethylene the carbon uses a different hybrid to form the s-bonds to H and C. The hybridization is determined (at least partially) by the electronegativity of the group, which is Bent’s Rule, defined by Henry Bent in 1961.20 Bent’s Rule states that hybrids used to bond to more electronegative substituents have higher p-character. The reason for this can be illustrated by considering a very common situation: an atom with a hybrid containing a lone pair and a heterocovalent 2-center 2-electron bond (Fig. 10). A lone pair and a s-bond both contain a pair of electrons. However, the bonding electrons are stabilized by two different nuclei (E and X); as a result, the central atom E can use a higher energy hybrid with more p-character to support the electron density in the bond. In the lone pair, the electron density must be supported only by the central atom E, so it uses an s-rich low energy hybrid. The same arguments can be applied to systems with two different bonds, such as E(Br)(Cl). Since chlorine is more electronegative and gives additional aid in supporting the bonding electron density, the central atom E uses a higher energy, p-rich orbital for the bond to chlorine and a more s-rich hybrid to bond to bromine. As discussed earlier, the angle between hybrids comes about so that orthogonality is maintained. By Bent’s rule, the type of hybrid used is based on the substituent. As a result, the structural angles in a main group compound are closely related to the substituents. If we know the structure of a compound, in many cases, we can define the hybrids being used by the central atom to form the bonds. As mentioned at the beginning, the structure is of great importance; the input for a Hybridization theory or Molecular Orbital theory analysis is always the structure. As a simple example, consider the structure of methyldichloroborane, MeBCl2. The compound exhibits Cs symmetry, with a Cl–B–Cl angle of 117.4 . By knowing no other information than this one angle, Hybridization theory can be used to calculate what hybrids are being used on boron to form all three bonds.

First, there are two hybrids that should be the same, the two used to bind to the chlorines. These should obey Eq. (18), and the hybrid used on boron to bond to the chlorines (lBCl) can be simply calculated as lBCl ¼ 2.2. This can be compared with the value obtained from NBO (B3LYP/6-311 ++G(d,p)) for lBCl(NBO) ¼ 2.35. cos oi¼j ¼ cos 117:2 ¼ −

1 lBCl

lBCl ¼ 2:2 In MeBCl2 there is an empty (except for possible weak p-interactions) p-orbital; as a result, there is one “hybrid” on boron of lpz ¼ 1. One can then use the Sum Rule (Eq. 15) to determine the only remaining boron-based hybrid, the hybrid used to form the B–C s-bond (lBC). The value obtained using this by-hand method is lBC ¼ 1.7, while the quantum calculation gives lBC(NBO) ¼ 1.5. X i

1 1 1 1 ¼2 ¼1 + + 1 + li 1 + lBCl 1 + l pz 1 + lBC ¼2



 1 1 1 + + ¼1 1 + 2:2 1 + 1 1 + lsBC lsBC ¼ 1:7

Since orthogonality is dependent upon the simple equations above, one may wonder why the values calculated by hand are not precisely the same as the values from the NBO calculation. The most common cause of the small deviations is “bond bending.” In the example above, it was assumed that the position of the group on boron is the exact direction that the hybrid points, which is

Fig. 10 Example illustrating Bent’s Rule. Both the bond to X and the lone pair contain a pair of electrons, but the bonding pair can use a higher energy hybrid (more p-character) because the electron density is also supported by group X. The lone pair, which is only supported by the central atom E, will use a lower energy s-character hybrid to stabilize the electron density.

Models for Understanding Main Group and Transition Metal Bonding

17

often a good approximation, but there are instances where this assumption is not accurate. Bond bending is a deviation of the hybrid used in a bond from the internuclear vector. To illustrate, consider cyclopropane. This small ring is described in organic classes as having “ring strain,” a manifestation of bond bending. The expectation is that the 60 angle in the triangular molecule must lead to bond bending as the closest possible approach of two orbitals in the main group is 90 with two pure p-orbitals. In actuality, the bond bending is significantly higher than might be expected, which can be determined from the geometry of the compound. Electron diffraction experiments on cyclopropane show that the structure has an H–C–H angle ¼ 114.5(9) .21 Assuming there is no bond bending in the C–H bonds, one can calculate using the Coulson Directionality theorem the carbon hybrid used to bond to the hydrogens as lCH ¼ 2.41. cos o ¼ −

1 lCH

cos 114:5 ¼ −

1 lCH

lCH ¼ 2:41 One can then find the hybridization parameter associated with the C–C bonds using the Sum Rule (Eq. 15). X i

2 2

1 ¼1 1 + li

1 1 +2 ¼1 1 + lCH 1 + lCC

1 1 +2 ¼1 1 + 2:41 1 + lCC lCC ¼ 3:84

As shown lCC ¼ 3.84; in other words, the bonding is being done by a hybrid quite far from a pure p-orbital, and the angle between hybrids will be larger than 90 . To find the angle, one can again use the Coulson Directionality theorem, although we are using the equation in reverse from the previous examples to find the angle between the hybrids. cos oCC ¼ −

1 lCC

cos oCC ¼ −

1 3:84

oCC ¼ 105:1 To find the bond bending angle (y) then just becomes a simple geometry problem (Fig. 11). The angle between the hybrids is 105.1 , and the C–C–C internuclear angle is 60 . The difference between these two is then 105.1 − 60 ¼ 45.1 . The bond bending is how far one of the hybrids is from the internuclear vector, which is half of 45.1 or 23 . NBO analysis gives bond bending angles as part of the typical output, and the bond bending angle in cyclopropane from NBO is 24 , almost identical to this back-ofthe-envelope calculation. Bond bending can occur under a variety of different circumstances, including geometric restrictions like ring strain and in many cases where hyperconjugation occurs in the molecule.

Fig. 11 (Left) Geometry of cyclopropane. (Right) Analysis of the bond bending.

18

1.02.6

Models for Understanding Main Group and Transition Metal Bonding

Lone pair bond weakening and its influence on organometallic chemistry

The reason there are three different values for covalent bond energies of the elements in Table 1 is due to “Lone Pair Bond Weakening” (LPBW), which is related to the primogenic effect and Bent’s Rule (vide supra). As it sounds, LPBW is a weakening of a bond from the Polar Covalence theory expected value due to the presence of a lone pair. The cause is best illustrated with an example, and we will consider ammonia.15 The experimental bond energy, mean enthalpy required to break all the N–H bonds, of NH3 is 93.4 kcal/mol,12 and Polar Covalence theory reproduces this number almost exactly (94.1 kcal/mol), if one uses the “fully weakened” value for the N–N bond energy (E0 ¼ 38.8 kcal/mol). According to Polar Covalence theory, if there were no LPBW, the bond energy in NH3 would be 128 kcal/mol, an enormous 34 kcal/mol stronger! What causes this large effect on the bond energy? To see LPBW in action, let’s examine what happens when we stretch one of the bonds in NH3, essentially looking at the process associated with measuring the Bond Dissociation Enthalpy. (This is somewhat different from the mean bond energy discussed in the previous paragraph, which is the atomization energy divided by 3.)e As one of the N–H bonds is stretched, the energy of the system rises until it reaches DH ¼ 106.7 kcal/mol (G4 energy using Gaussian16), the BDE for the first bond of ammonia (Fig. 12A). As shown in Fig. 12B, as the bond is stretched, the nitrogen hybrids for the unstretched N–H bonds change relatively little from sp2.8 in the ground state to sp3.5 in NH2, but these do lose s-character as the H–N–H angle changes from 107.9 in NH3 to 101.9 in NH2. In contrast, the hybrid parameter used by nitrogen to bond to the hydrogen being stretched increases exponentially (Fig. 12C) from sp2.8 to sp1, a pure p-orbital. As all the nitrogen hybrids used for bonding to the conserved hydrogens lose s-character, the lone pair hybrid orbital gains significant s-character, changing from a hybridization of sp3.6 to sp0.8. A summary of some of these effects is shown at the top of Fig. 12, i.e., the lone pair obtains additional s character and the orbital previously used to bond to hydrogen becomes a pure p orbital holding a radical. Since the s-orbital is lower in energy, it makes stronger bonds with substituents attached to nitrogen. Based on the calculations, one can surmise that the LPBW effect in NH3 is due to s-orbital “stealing” by the lone pair, which significantly lowers the N–H BDE. Again, Bent’s rule suggests that the lone pair will be held in an s-rich orbital, it is essentially this fact that leads to N–H bond weakening in ammonia as the bonds must then be formed with higher energy p-rich hybrids. The LPBW effect is expected to be most significant for 1st row p-block elements as the primogenic effect leads to the 2s and 2p orbitals being very nearly the same size, which in turn leads to extensive hybridization. Since the heavier p-block elements do not hybridize as well, there is less s character to “steal” from the bonding orbitals, which are already predominantly composed of p-orbitals. This leads to many of the familiar differences between the 1st row and heavier elements, such as weak O–O single bonds, weak N–N single bonds, N2 being the preferred state of elemental nitrogen, P4 (and other s-bond structures) being preferred elemental forms for phosphorus, and many more. LPBW effects are alleviated by increasing the number of p-bonds to the atoms—since p-bonds by necessity only use p orbitals they free up s character for the remaining s bonds and lone pairs. The fully weakened covalent bond values (E0 ) are used for single bonds, partially weakened (E00 ) for double bonds, and fully unweakened (E000 ) for triple bonds. In addition, adding a Lewis acid to the lone pair can alleviate LPBW as the resulting dative bond requires less s contribution from the atom than the lone pair. There are many examples where LPBW is pertinent to organometallic chemistry, and one that will be briefly mentioned here is the bonding of carbon monoxide. CO has the strongest bond known (256 kcal/mol). The bond is quite short at RCO ¼ 1.128 A˚ with a stretching frequency of nCO ¼ 2143 cm−1. Polar Covalence theory can be used to calculate a bond energy that is close to this value when using Sanderson’s “fully unweakened” bond energies, as expected for a triple bond (vide supra). The bonding of CO to a metal center is likely familiar to anyone who has had an inorganic or organometallic chemistry course (Fig. 13). A metal center with an empty acceptor orbital of s-symmetry can act as a Lewis acid toward the carbon-based HOMO, a lone pair. In addition, metal centers with a filled d-orbital of p-symmetry can “backdonate” electron density into the LUMO of CO, the C–O p-antibonding orbitals. One often discussed consequence of CO binding to a metal center where backbonding is possible is a dramatic decrease in C–O stretching frequency due to occupation of the p-antibonding orbitals on the ligand. For example, the IR-active T1u stretch in Cr(CO)6 is found at 2003 cm−1, over 100 cm−1 lower than the free CO stretch due to this backbonding.22 Less discussed is the effect of the s-lone pair interaction of CO with the metal center, which is most evident when CO is bound to a metal without p-backbonding. In these cases, metal binding leads to a stronger C–O bond with a shorter C–O distance and a higher stretching frequency. Complexes with these properties of having higher C–O stretching frequencies than free CO are what have been called “nonclassical carbonyls.” From an MO perspective, the lone pair has some s-antibonding character due to configuration interaction with an antibonding orbital of the same symmetry. In the Hybridization theory parlance used thus far in this writing, the lone pair in free CO has more s-orbital character than the C–O s-bond due to Bent’s Rule. When CO binds to a metal center, the lone pair loses some of the “extra” s-character and the C–O s-bond gains additional s-character, making the bond stronger. In other words, CO has LPBW that is relieved by coordination to a Lewis acid. An example of a nonclassical carbonyl is Ag(CO)+, which has a stretching frequency of 2204 cm−1, 50 cm−1 higher than free CO! NBO calculations suggest that free CO has orbital hybridizations for the lone pair on carbon and the carbon hybrid used to form the e NBO analysis of NH3 gives a somewhat surprising outcome—there is a small amount of bond bending, which causes the compound to disobey Bent’s Rule. Consequently, the lone pair has slightly more p-character than the hybrids on nitrogen used to form the N–H bonds, but this is a side issue that doesn’t significantly affect our current discussion of LPBW.

Models for Understanding Main Group and Transition Metal Bonding

a

19

b

c Fig. 12 (Top) Hybrids found in NH3 and NH2 + H. (A) Energy of NH3 as one of the N–H bonds is stretched until broken, i.e., the BDE of the first bond of ammonia. (B) NBO hybridizations of the nitrogen orbitals used to bond to the hydrogens not being stretched and the hybridization of the lone pair-containing orbital. (C) NBO hybridizations of the nitrogen orbitals bonding to H in NH3 as one of the bonds is stretched.

s-bond of sp0.3 and sp3.1, respectively. When Ag+ coordinates to CO, the lone pair hybridization used to bond to silver decreases in s-character to sp0.5, and the carbon then uses an sp1.8-hybrid to form the s-bond to oxygen. This C–O bond strengthening due to alleviation of Lone Pair Bond Weakening is always present when CO bonds to a metal center and is a function of the metal center’s Lewis acidity. The strengthening is expected to run counter to the typically observed weakening of C–O bond strength due to occupation of the p-antibonding orbitals, when present. As mentioned, LPBW is likely an important effect for many complexes. For example, first row p-block elements with lone pairs, e.g., fluoride, alkoxide, and amide, are more common ligands for high oxidation state metals (often early metals) with empty

20

Models for Understanding Main Group and Transition Metal Bonding

Fig. 13 Bonding of CO to a metal center.

p-acceptor orbitals where LPBW can be alleviated by interaction with the metal center. Amides, for example, are certainly known for later, low-oxidation state metals with groups such as aryl and silicon on the nitrogen to help stabilize the lone pair. Late metal fluoride complexes, such as the fluoride derivative of Wilkinson’s catalyst FRh(PPh3)3, often show unusual reactivity because the metal has no empty p-acceptor orbitals to alleviate LPBW.23,24

1.02.7

Beyond the 2-center 2-electron bond

The focus above has been on 2-center 2-electron (2c2e) bonds, but these are by no means the only types of bonds found for main group or transition metal compounds. Many chemical species, however, can be understood as a combination of 2c2e-bonds and just three other bonding types, which will be briefly discussed here. It was once thought that there was d-orbital participation in the main group to explain the stability of hypercoordinate compounds like XeF2 and OPMe3, but this is unnecessary, and contradicted by quantum chemical calculations. Most of the chemical community at this stage has abandoned the idea of significant d-orbital participation in the main group.19,25 Instead, the various observations in the main group can be explained using models in which the number of electrons around the central atom does not exceed an octet, because the bonds are not 2-center 2-electron bonds. Instead, they use other bond types, and these same bonding interactions will be discussed for transition metals as well. The 3-center 4-electron bond (3c4e or o-bond) often is found in cases where there are too many s-bonds in a Lewis structure for the number of possible hybrids on the central atom. One of the simplest examples of such a compound is XeF2. The free Xe atom, being a noble gas, has eight valence electrons. The typical method for drawing the Lewis structure is shown at the top of Fig. 14, where there are three LP on each atom and a bond to each fluorine; however, this suggests that Xe is sharing 10 valence electrons, despite the presence of only four valence orbitals (one s- and three p-orbitals). The molecular orbital explanation for the bonding in XeF2 can be drawn very simply (Fig. 14), but first it will be useful to borrow a little from Hybridization theory, namely Bent’s Rule. As mentioned, there are three lone pairs on each atom of XeF2. Bent’s rule

Fig. 14 The typical Lewis (top), partial MO (middle), and hybridization (bottom) descriptions for XeF2.

Models for Understanding Main Group and Transition Metal Bonding

21

states that lone pairs prefer to be in s-rich orbitals, leaving the bonding orbitals higher in p-character (Fig. 10). Since there are so many lone pairs on all the atoms, we will make the simplifying (and very good) assumption that the bonding is being done by pure p-orbitals. As a result, there is only one p-orbital remaining on each of the two fluorines and xenon to form the bonds. The three p-orbitals on the three atoms will form MOs with 0, 1, and 2 nodes (between atoms) that are bonding, nonbonding, and antibonding, respectively. These three MOs are filled with four electrons. The valence electrons in the system are 8 (Xe) + 7 (F) + 7 (F) ¼ 22 electrons, minus the electrons in the three lone pairs on each atom, i.e., 3  3  2 ¼ 18 electrons in lone pairs. So, the electrons in our MOs are 22 − 18 ¼ 4 electrons, which fills the bonding orbital that spans the entire molecule and the nonbonding orbital based on the fluorines, the two MOs lowest in energy. In other words, there is a bond to Xe that is split between the two fluorines for a bond order of ½ to each fluorine. The valence bond description is also shown in Fig. 14, which has a full bond between one Xe and F and no bond to the other F in one resonance form. This contributes equally with a resonance form where there is a bond to the opposite fluorine. As a result, there is a ½ bond between each F and Xe in total. It is precisely these two and only these two resonance forms, with 50:50 contribution, that are found when the system is investigated computationally using Natural Resonance Theory (B3PW91/SDD). When a system contains what appears to be too many s-bonds for the number of possible hybrids, these o-bonds (3c4e-bonds) are often present. Another common type of bonding interaction found when compounds have Lewis structures that seem to have too many p-bonds (causing there to be more than eight electrons) to a central atom is negative hyperconjugation. Negative hyperconjugation can be defined as “donation from a filled nonbonding orbital into an empty antibonding orbital,” which is quite analogous to the backbonding discussed earlier between a metal and CO (Fig. 13). A relatively simple form of negative hyperconjugation occurs when phosphines bond to some metal centers. If a metal center contains unpaired electron density, e.g., a lone pair in a d-orbital, that electron density can donate into the s-antibonding orbitals, e.g., P–C s -orbitals on the phosphine (Fig. 15). The expected consequence is a relatively short M–P bond as the bond order here is increased and a relatively long P–C bond as the antibonding orbital is occupied. One of the many examples is shown in Fig. 15. In the iron complex on the left the metal interacts with an allyl-like group and three phosphite ligands. The iron complex on the right is identical except for being oxidized by one electron (a BF−4 counterion is not shown). Upon oxidation, the metal center’s ionic radius should decrease, and this can be seen in the Fe–C distances, which are significantly shorter in the oxidized complex. However, the Fe–P distances increase when the complex is oxidized, which is explicable as the oxidized metal having less capability for “backbonding” (negative hyperconjugation) to the phosphites. Also, as expected, the cationic complex with less negative hyperconjugation has shorter P–O distances due to lesser occupation of the P–O antibonding orbitals on the phosphite. The backbonding is increased by having a more electron-rich metal center and a phosphine-derivative (PX3) with substituents X that can support a negative charge well. At the extreme, PF3 is estimated to be as good or better a p-electron withdrawing group than CO.28 Negative hyperconjugation appears in innumerable metal complexes and main group compounds. It is often most noticeable in cases where a compound has more hybrids than would seem to be available because of p-bonded ligands. An example from the

Fig. 15 Structures of the redox isomers {Fe(Z3-C8H13)(P(OMe)3)3}0/+1. The higher oxidation state complex shows less negative hyperconjugation into the phosphites.26,27

22

Models for Understanding Main Group and Transition Metal Bonding

main group is a phosphine oxide, often drawn as O]PX3 implying a 10-electron central phosphorus atom in the limit of 2-center 2-electron bonds. Typically, the most prevalent resonance form is zwitterionic (and octet rule compliant) − O–P+ X3 with negative hyperconjugation between the oxygen lone pairs and P–X s -orbitals. Outside this example, negative hyperconjugation is critical to understanding the bonding in sulfuric acid and phosphoric acid, some of the largest commodity chemicals in the world. The third common bonding type outside of the prototypical 2c2e bond to be covered here is the 3-center 2-electron bond (3c2e, t-bond). While 3-center 4-electron bonds (o-bonds) and negative hyperconjugation are prevalent in hypercoordinate compounds, the 3c2e bond is commonly found in hypovalent compounds. One of the simplest examples of this bond type is in the ethyl cation (or protonated ethylene), which has a non-classical structure with a hydrogen bridging between the two carbons in the gas phase.29–31 In attempting to draw a composite structure for protonated ethylene, the Lewis structure at the top of Fig. 16 is undesirable because the central H exceeds the available valence if 2c2e bonds are used. From an MO perspective, we can consider what would happen if a proton were to interact with the HOMO of ethylene, the p-bonding orbital. The HOMO of ethylene would be transformed by interaction with the proton to give 3-center 2-electron bonding and antibonding orbitals (t- and t -orbitals), while the LUMO would be of the wrong symmetry to interact with the incoming proton and would be more or less the same as in ethylene. The Hybridization theory picture, via Natural Resonance theory, is shown at the bottom of Fig. 16, which has about 50% of the ground state comprised of ethylene with a noninteracting proton and 2  25% with a bond between the bridging hydrogen and either of the two carbons. Prototypical examples of t-bonds are found in the AlMe3 dimer, which exhibits C2h symmetry by neutron diffraction (Chart 1).32 Main group and early transition metals in high oxidation states display 3c2e bonds most often, but there are many examples involving later transition metals as well like the simple {(m-Me)Rh(COD)}2 and the beautiful Cp Cr(m-CH2)(m-CH3)2CrCp .33,34 Please note that bridging halides are not considered to have t-bonds, because they have multiple available orbitals for 2c2e bonds. A bridging halide is most often interpreted as having a covalent bond to one metal and a dative bond (from one of its lone

Fig. 16 (Top) A line drawing of the protonated ethylene (ethyl cation) with a bridging hydrogen. Since the hydrogen has only one valence orbital, the two lines cannot represent 2c2e bonds, but represent at 3c2e bond (t-bond). (Middle) The MO diagram of protonated ethylene showing the t- and t -orbitals. (Bottom) Results of a NRT calculation of protonated ethylene showing the major contributing resonance forms. (Some minor contributors are not shown.)

Models for Understanding Main Group and Transition Metal Bonding

23

Chart 1 Examples of stable compounds with t-bonds.32–34

pairs) to the other metal, with several resonance forms possible. Likewise, bridging halides, alkoxides, amides, etc. can form two 2c2e-bonds rather than use a t-bond. In a t-bond, two electrons are shared by three orbitals across three different atoms. This type of sharing of electron pairs across orbitals is common in organometallic chemistry. There are many related bond types where >3 orbitals share a pair of electrons. For example, tert-butyllithium exists as a tetrahedron of lithium atoms with bridging tert-butyl groups on the faces, implying carbons involved in 4-center 2-electron bonds with three lithium atoms. In addition to these relatively common bonding types, there are also 2-center 3-electron bonds in many molecules, including simple diatomics like nitric oxide and dioxygen. Nevertheless, the three bonding interactions discussed—3c4e-bonds (o-bonds), negative hyperconjugation, and 3c2e-bonds (t-bonds)—are sufficient to enable a cogent discussion of the bonding of most complexes with interactions other than 2c2e bonds.

1.02.8

Examples of using Bent’s rule, hybridization, and structure in the main group

The ideas discussed above can be used to understand the structures of main group compounds, with an eye toward a description of electronic structure. While VSEPR theory provides a quick method for guessing the structure of compounds from their formulas, ideas like Bent’s rule, Hybridization theory, and the various bonding types provided in previous sections can be at least as (and often more) useful in deducing structure. In this brief section, some examples of analyzing main group bonding “by hand” will be given with examples of 3c4e-bonding, negative hyperconjugation, and 3c2e bonding. An example that illustrates several of the principles above is determination of the electronic structure of ClF3, a very strong oxidant. The common Lewis structure of the molecule gives two lone pairs on the central chlorine atom with three bonds to fluorines, which suggests 10 electrons around the central atom. As is obvious from the electron count over eight for the main group compound, this structure is not a valid resonance form, and the high electron count is due to “extra” s-bonds, a situation that often suggests that the compound contains one or more 3c4e-bonds (o-bonds). In this case, one would expect only one o-bond, which would bring the electron count of Cl to 8. In an o-bond, the two groups are typically 180 apart, which gives one anticipated structure—T-shaped. The molecule is, in fact, approximately T-shaped as expected from this analysis (also consistent with VSEPR). The geometry of ClF3 from a DFT optimization (B3LYP/6-311 ++G ) is shown in Fig. 17. The two fluorines involved in the 3c4e-bond are 175 apart, as expected when they mutually interact with an approximately pure p-orbital on Cl. These o-bonded fluorines have a lower bond order, which is confirmed by longer Cl–F distances, calculated as 1.77 A˚ . About 90 from the o-bond axis is a fluorine better described as being involved in a 2c2e-bond with a distance of 1.66 A˚ . The chlorine also has two lone pairs, and, consistent with Bent’s Rule, a large portion of the s-orbital on chlorine is used to stabilize these lone pairs, along with a p-orbital. NBO and NRT calculations are fully consistent with this picture of ClF3, and some of the results are shown in Fig. 17. The NRT suggests that only three resonance forms contribute—two resonance forms consistent with the o-bond description that together contribute 80% of the ground state structure and a third resonance form where there are covalent bonds to the two fluorines 180 apart (where Cl uses sp-hybridized orbitals) and no bond to the third fluorine. The slight amount of bond bending in the compound making the angles slightly different from 180 and 90 is attributable to VSEPR-like steric interactions between the lone-pairs and bonds.19 The simplest isolable phosphorus ylide is often drawn as H2C]P(CH3)3, which was prepared and isolated by Schmidbaur and Tronich in 1968 (Fig. 18).35 Again, this example has more than eight electrons around the central atom in the usual line drawing. The solution to this was given in the description of the compound as an “ylide,” H2C−–P+(CH3)3, and it is this resonance form that is expected to be the major contributor. However, some negative hyperconjugation is also expected, which increases the P–CH2 bond order somewhat. In a molecule like OPCl3, the axially symmetric oxygen donates symmetrically into all the antibonding orbitals of the chlorides. In contrast for H2CPMe3, the plane of the CH2 lone pair is directed at one specific methyl group acceptor; as a result, one P–Me bond is expected to be longer than the others. NRT (NBO6, B3LYP/6-311 ++G ) calculations suggest that the main contributor is indeed the ylide structure at 62%. The second largest contributor, 19%, is due to negative hyperconjugation where the unique methyl-P antibonding orbital acts as the acceptor, giving a structure with an H2C]P double bond and no bond from phosphorus to one of the methyl groups.

24

Fig. 17 of ClF3.

Models for Understanding Main Group and Transition Metal Bonding

(Top) The geometry of ClF3 from a DFT (B3LYP/6-311++G ) optimization. (Middle) Hybrids from the NBO analysis of the molecule. (Top) NRT analysis

Fig. 18 (Top) Geometry of H2C]PMe3 from DFT (B3LYP/6-311++G ) optimization. (Middle) An MO view of negative hyperconjugation between the ylide carbon-based lone pair and the P–Me antibonding orbital. (Bottom) The major contributing resonance forms from the NRT analysis.

As a final example, consider the structure of the AlMe3 dimer with its t-bonded bridging methyl groups (Fig. 19). Since the t-bonds have three nuclei (two Al and one C) to support only two electrons, Bent’s rule suggests that the aluminum could use p-rich hybrids for these bonds and more s-rich hybrids for the 2c2e-bonds to the terminal methyl groups. As a result, the Meb-Al–Meb (Meb ¼ bridging methyl) angle is expected to be closer to 90 than the Met–Al–Met angle (Met ¼ terminal methyl). This is precisely what is found experimentally, with an Meb-Al–Meb angle of 102.0(3) and an Met–Al–Met angle of 125.8(3) .32 The nature of the hybrid used by aluminum to bond to the bridging methyl is a little less obvious than usual as one would expect that the hybrid is directed a little away from the carbon in a t-bond to concomitantly interact with the other aluminum orbital as well. As a result, one expects the two Al hybrids used in the t-bonds to be at an angle slightly less than the Cb–Al–Cb angle. The experimental angle between the two terminal methyl groups on an aluminum atom in the AlMe3 dimer is 125.8 , and applying the Coulson Directionality Theorem to this angle between two equivalent hybrids indicates that aluminum uses an sp1.7-hybrid to form these bonds. We can find the t-bonding hybrids (lt) by employing the Sum Rule (Eq. 15) and the hybridization used to bond to the terminal methyl groups (sp1.7). The result is shown below, which is that the aluminum uses very p-rich sp6.6-hybrids to bond to the bridging methyl. The equivalent t-hybrids on aluminum are then (Coulson Directionality theorem) 98.7 apart, close but slightly more acute than the experimental Meb-Al–Meb angle of 102.0(3) .

Models for Understanding Main Group and Transition Metal Bonding

1 2

1

1 1 1.7

25

1 2

1

1

6.63

Fig. 19 The angle between the hybrids used to bond to the bridging methyl groups is slightly more acute (98.7 ) than the experimental angle between the nuclei (102.0 ). (Angle difference shown is a bit exaggerated to make it noticeable.) 4 X i¼1

    1 1 1 +2 ¼2 ¼1 1 + li 1 + 1:7 1 + lt lt ¼ 6:63

1.02.9

Hybridization theory and the transition metals

Hybridization theory for the transition metals is more complex than for the main group, and development of hybrid models lags behind for the transition metals relative to the p-block. Further, MO theory and Ligand Field theory are far more developed at this stage. Nevertheless, there are some hybridization models for transition metal system, and they have proven quite useful in some cases. For example, Hoffman’s Isolobal Analogy is a hybridization model based on d3sp2 hybrids.36 In addition, Landis and coworkers have developed a model for understanding transition metal bonding using sd-hybridization, arguing that the p-orbitals of these metals are high in energy and participate very little.37–39 For d-orbital hybridization, one can discuss a new hybridization parameter, m, and sdm-hybrids. Just like for spl-hybrids, a specific angle between the hybrids is a requirement for their orthogonality, and an analog of the Coulson Directionality theorem relating the angle (oij) between two hybrids (mi and mj) to the hybridization parameters is shown in Eqs. (20a) and (20b).19 sffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi 2 ffi 1 − pffiffiffiffiffiffi mi m j when m > 2 (20a) cos oij ¼  3 cos oij ¼ 0 when m 2

(20b)

One can understand the nodal relationship by considering the consequence of mixing a dz2-like orbital with an s-orbital (Fig. 20). In essence, as the amount of s-character in the sdm-hybrids increases, the size of the torus will decrease changing the angular location of the node. A pure d-orbital has nodes at 125 and 55 . As the amount of s-orbital mixing increases (lower m) the angles of the nodes both tend toward 90 , i.e., the angle of the 55 node increases and the angle of the 125 node decreases until they become equal at 90 . This extinguishing of the torus and a 90 angle for orthogonality is reached when m 2 (Eq. 20b). In determining what sd hybridization is applicable to a particular system, it is important to keep orbital symmetry and Bent’s Rule in mind. If the complex has unpaired electrons on the metal center, based on Bent’s Rule, the lowest energy orbital available will be used in large part to house these electrons; in other words, lone pairs on the metal center will reside in approximately pure d orbitals. If there are p bonds to the metal center, since we are ignoring p-orbitals, pure d-orbitals must be used to form these p bonds. Using these hybridizations and their preferred angles, one can deduce the preferred geometries for homoleptic complexes containing these hybridizations for normal and hypovalent complexes. Examples for sd1 to sd5 are shown in Fig. 21. Under the sd5 to sd1 hybridizations are shown the geometry that is likely to be present using the preferred angles and the idealized point group of such a structure. Then, there is an example of a compound with the sdm hybridization. Just as in the main group, many transition metal complexes are hypercoordinate, and the same types of bonding models we used previously are useful here. Compounds that are hypercoordinate due to s-bonds often have 3c4e-bonds (o-bonds) present with the bonding type extending along an axis at 180 . If the complex is hypercoordinate with p-bonds, then there are often negative hyperconjugation interactions present.

26

Models for Understanding Main Group and Transition Metal Bonding

55° 125°

63°

71°

117°

109°

sd5-hybrid

90° sd2-hybrid

sd3-hybrid

(as dz2)

hybridizaƟon parameter (m) ≤2 3 4 5 7 10 12 20

Small Nodal Angle 90 70.5 65.9 63.4 60.8 58.9 58.2 56.8

Large Nodal Angle — 109.5 114.1 116.6 119.2 121.1 121.8 123.2

Fig. 20 As the amount of s-orbital mixed with dz2 increases, the torus size decreases changing the angle between the main lobe of the orbital and the node. Once m 2 in the sdm-hybrid, the preferred angle between hybrids becomes 90 . This angle required for hybrid orthogonality can be calculated using Eqs. (20a) and (20b), and the angle between identical hybrids is shown in the table in the figure for a few select hybridizations.

sd5

W

sd3

Me

Me Me

Me Me

sd4

Me Me

Me Me

Ta

Ph Os

Me Me

Ph

sd2

Ir mes

Ph Ph

sd1

Ar2B mes mes

Cr N

N

BAr2

Ar Ar Ar = mesityl

Fig. 21 Examples of homoleptic complexes with sd1–5 hybridization and the idealized geometries of the hybridizations. In some cases, the example compounds show some bond bending, and the ideal angles between hybrids does not correlate with ligand-metal-ligand angles.

A very common system type that is hypercoordinate under this model is square planar, d8-complexes, which are common in Rh(I), Ir(I), Pd(II), Pt(II), and other systems. Here, we use a Pt(II) square planar complex as an example, {Pt(NH3)4}2+. As would be expected, NBO analysis (NBO6 on density from B3LPW91/SDD) places four lone pairs on the platinum center, all in approximately pure d-orbitals. Consequently, the highest hybrids possible for the metal to use in bonding to the nitrogens are sd-hybrids, which are 90 apart (vide supra). The NBO analysis also gives the hybridization of the orbitals used by Pt to bond to N as sd1.1. One such orbital is shown at the bottom of Fig. 22. An NRT calculation gives four major resonance contributors that participate equally. These

Models for Understanding Main Group and Transition Metal Bonding

2+

NH3 H3N

Pt

27

NH3

NH3 NH3 H3N

Pt NH3

NH3 NH3

H3N

Pt

NH3 NH3

H3N

NH3

Pt NH3

NH3 NH3

H3N

Pt

NH3

NH3

Fig. 22 (Top) Line drawing of {Pt(NH3)4}2+. (Middle) NRT analysis of the complex gives these four resonance forms, all of which contribute 25%. (Bottom) Drawing of one of the bonding NBOs between the platinum and ammine nitrogens.

have bonds to the nitrogens 90 apart owing to the Pt sd-hybrids used. The ammine ligand 180 from each Pt–N bond is donating into the antibonding orbital of the s-bond across from it; in other words, these are 3c4e-bonds, like those in XeF2, involving o bonds from sd-hybrids instead of the pure p-orbital used in the Xe case. Like the main group examples discussed, the central transition metal will predominantly use 2c2e-bonds when it has enough hybrids to do so, but if there are too few hybrids available (e.g., due to unpaired electron density on the metal or p-bonds that require d-orbitals) 3c4e bonds and negative hyperconjugation interactions are common.

1.02.10

Practical evaluation of metal-ligand bond interactions for applications in catalysis

The ideas above, with their attendant approximations, are useful in understanding the electronic structures of main group and transition metal organometallic compounds. However, they can have great impact: for example, the trillion-dollar industry of catalysts and their products depends upon understanding metal-ligand interactions and their effects on reactivity. In addition, quantum mechanical calculations can lead to refinement of these ideas and improved catalysts. While much of the above discussion was qualitative, quantitative methods for evaluating ligand donor properties and sterics can be used to model reactions for catalyst optimization. The results of such studies often have led to a better understanding of ligand effects in a particular reaction.40–46 For example, Fey and coworkers have developed a phosphine “map” for applications in catalysis.47 The stereoelectronic parameters of the map, which are largely derived from quantum chemical calculations, have been used to improve ee’s in Tsugi-Trost allylations48 and to examine ligand effects in individual steps of Suzuki couplings.49 The prototypical example of such ligand modeling was provided by Tolman for phosphine-containing complexes.50 Tolman used the symmetric CO-stretching frequency from infrared spectroscopy in LNi(CO)3 complexes as his quantitative measure of the donor ability of the phosphine, L.50 As mentioned above, as the metal center becomes more electron-rich, it can backdonate more strongly into the CO ligands, reducing their stretching frequencies through negative hyperconjugation. The values are finally rescaled so that the frequency measure for L ¼ PtBu3, a very strongly donating phosphine, is 0 cm−1. This electronic parameter (w) was combined with a steric parameter, the Tolman Cone Angle (y), to model reactions involving low oxidation state metals with phosphine ancillary ligands. Some values are given in Fig. 23.51 The Tolman Cone Angle was found using physical models with a special ruler built to determine the angle between van der Waals radii of atoms on one side of the ligand to the other when bound to a metal; larger ligands take up a larger angle of space near the metal center.50 To understand some property (Z), e.g., reaction rate, stereochemistry of a product mixture, spectroscopic measurement, etc., of a low oxidation state system with phosphine ancillary ligands, Tolman used a simple model involving the Tolman Electronic Parameter (w) from infrared spectroscopy on LNi(CO)3 and the Tolman Cone Angle (y). In other words, the minimalist model (Eq. 21) would use one parameter for the electronics (w) and one for the sterics (y), where a, b, and c are coefficients found by least squares fit.50 Z ¼ a + bðwÞ + cðyÞ

(21)

A limitation of this method is that it is applicable directly only to phosphines and closely related ligands (e.g., NHC). More recently, analogous methods have been developed for high oxidation state metals where a wider array of ligand types are prevalent. One way of quantifying the nature of an X ligand uses a chromium(VI) nitride, NCr(NiPr2)2X, as the reference experimental system for

28

Models for Understanding Main Group and Transition Metal Bonding

Fig. 23 Some representative values for the Tolman Electronic Parameter used for low oxidation state metal complexes and Ligand Donor Parameter for high oxidation state complexes.

determination of donor ability. In this case, the diisopropylamides have some multiple bond character to the chromium center due to the dative interaction between the nitrogen lone pair and the acceptor orbitals on the metal. However, as X becomes a better donor (both s and p) it competes with the lone pair on the amide for the metal’s acceptor orbitals. As X becomes a better donor, there is less p-bonding to the amides due to the competition, and the amides rotate more rapidly. The rate of rotation of the diisopropylamides in this system are readily measured using spin saturation transfer in the 1H NMR spectrum in toluene-d8. This rate constant can be converted to the barrier to rotation (DG{) using the Eyring Equation, and one assumes that the entropy associated with the rotation is invariant for NCr(NiPr2)2X, which allows calculation of the enthalpy of activation for the amide rotation, dubbed the Ligand Donor Parameter (LDP). Like the Tolman Electronic Parameter, more donating ligands have lower values. Some representative values are shown in Fig. 23.52–54 LDP values have been found to correlate with a wide range of properties for d0 metal complexes and have been used in a simple model to understand a catalytic system. The model employed was essentially that of Tolman shown in Eq. 21, where LDP was used for w and a different steric measurement (percent buried volume, %Vbur)f was used for y.53,54 The system modeled was simple hydroamination of an alkyne by a series of titanium catalysts with bidentate ligands.53 In this case all the chelates were symmetrical, with the same donor on both sides of the bidentate ligand. Once the model was established with 3-4 ligands being measured, most of the rates for the subsequent catalysts could be anticipated before the catalysts were even prepared. The model line and the fit are shown in Fig. 24.

kobs (x 10–4) = 1.34 + 1.61(LDP) – 2.25(%Vbur) Me Me

N N

Ti

NMe 2 NMe 2

6

6 Ph

F

6 6

N

H NMe 2

6 Ti

Me 2(H)N Me 2N

N NMe 2

F

Bu t

Ti

O

8

O

NMe 2 Bu t

7 7

Fig. 24 A plot of the calculated rate constant using the obtained model for a pseudo-first order hydroamination reaction vs the experimental rate constant. The scaled model (scaled so that the coefficients may be compared directly) is shown at the top of the figure. The numbers adjacent to the ring size of the ancillary ligand in the likely active species (see text). The points in blue are aryloxide-based catalysts. The points in red are indolyl- and pyrrolyl-based catalysts, and a few representative catalyst structures are shown.53

f Percent buried volume (%Vbur) approximates the primary coordination sphere around the metal then determines the percentage of that sphere occupied by the ligand(s) in question. The input is structural data, e.g., coordinates from X-ray diffraction.

Models for Understanding Main Group and Transition Metal Bonding

29

However, two catalysts did not fit the model, and it was found that they had different chemistry from the other catalysts. (One catalyst dimerized irreversibly, and another gave side reactions not seen in the other systems.) From the model, it was possible to gather some mechanistic information about the catalysis and predict the rate of proposed catalysts in the same class. For example, the model did nothing to take into account the ring size of the chelate; metallacycles with six to eight members were used with the ring size apparently not significantly affecting the reaction rate as only the donor ability and sterics were used in the modeling. In addition, some catalysts were structurally characterized as having Z5- and Z1-pyrrolyls, but all ligands were modeled as having the donor ability of an Z1-pyrrolyl, which suggests that the active species has only Z1-pyrrollides. Again, these are empirical results as all the data used in the model is from various experiments like NMR kinetics and X-ray diffraction (for %Vbur).53 The model from scaled parameters is shown at the top of Fig. 24, which allows the coefficients to be compared directly. The steric parameter (%Vbur) coefficient is negative and slightly larger than the electronic parameter (LDP) coefficient, suggesting that as the ligands become smaller the reaction rate increases and that this parameter is slightly more important than the electronics. The electronic coefficient associated with LDP is positive, suggesting that as LDP increases (more electron-deficient) the reaction rate increases. (The first coefficient is an intercept where the LDP and %Vbur are zero, which is not particularly meaningful but necessary for the fit.)53 The model with the natural (unscaled) parameters was used to predict the activity of new catalysts during the study. Consequently, these types of models can be valuable for reaction optimization, giving not only the changes one should make for better catalysts (in this case, more electron-deficient and smaller), but also quantitative expectations for the rates of novel catalyst possibilities where the supporting ligands have known parameters.

1.02.11

Concluding remarks

The above was a brief overview of some useful concepts to understand metal-ligand bonding, along with a cursory discussion of building quantitative models for catalysis. In some cases, these concepts help one to anticipate the important structural features of a compound, and in many cases the models employed can give insight into the bonding if one knows structure. In the end, from a pragmatic perspective, the “best” model for a particular system is the one that allows the practitioner to understand their data on some level and design the next experiment to move the science forward. In at least some cases, simple models and concepts like those discussed here can be powerful allies to discovery.

Acknowledgments The author thanks the National Science Foundation for support (CHE-1953254).

References 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22.

Pauling, L. The Place of Chemistry in the Integration of the Sciences. Main Curr. Mod. Thought 1950, 7, 108–111. Shaik, S.; Rzepa, H. S.; Hoffmann, R. One Molecule, Two Atoms, Three Views, Four Bonds?Angew. Chem. Int. Ed. 2013, 52 (10), 3020–3033. Zhao, L. L.; Pan, S.; Holzmann, N.; Schwerdtfeger, P.; Frenking, G. Chemical Bonding and Bonding Models of Main-Group Compounds. Chem. Rev. 2019, 119 (14), 8781–8845. Pyykko, P. Dirac-Fock One-Center Calculations. 8. 1-Sigma States of ScH, YH, LaH, AcH, TmH, LuH and LrH. Phys. Scr. 1979, 20 (5–6), 647–651. Kaupp, M. The Role of Radial Nodes of Atomic Orbitals for Chemical Bonding and the Periodic Table. J. Comput. Chem. 2007, 28 (1), 320–325. Schwerdtfeger, P.; Smits, O. R.; Pyykko, P. The Periodic Table and the Physics That Drives It. Nat. Rev. Chem. 2020, 4 (7), 359–380. Power, P. P. Bonding and Reactivity of Heavier Group 14 Element Alkyne Analogues. Organometallics 2007, 26 (18), 4362–4372. Waber, J. T.; Cromer, D. T. Orbital Radii of Atoms and Ions. J. Chem. Phys. 1965, 42 (12), 4116. Van Vleck, J. H. The Group Relation Between the Mulliken and Slater-Pauling Theories of Valence. J. Chem. Phys. 1935, 3 (12), 803–806. Shaik, S. S.; Hiberty, P. C. A Chemist’s Guide to Valence Bond Theory; Wiley-Interscience: Hoboken, NJ, 2008. Landis, C. R.; Weinhold, F. Comments on “Is It Time to Retire the Hybrid Atomic Orbital?”J. Chem. Educ. 2012, 89 (5), 570–572. Sanderson, R. T. Polar Covalence; Academic Press: New York, 1983. Sanderson, R. T. Electronegativity and Bonding of Transitional Elements. Inorg. Chem. 1986, 25 (19), 3518–3522. Herzberg, G.; Huber, K.-P. Molecular Spectra and Molecular Structure. IV. Constants of Diatomic Molecules; Van Nostrand: New York, 1979. Lauvergnat, D.; Maitre, P.; Hiberty, P. C.; Volatron, F. Valence Bond Analysis of the Lone Pair Bond Weakening Effect for the X-H Bonds in the Series XHn ¼ CH4, NH3, OH2, FH. J. Phys. Chem. 1996, 100 (16), 6463–6468. Sanderson, R. T. Electronegativities in Inorganic Chemistry. J. Chem. Educ. 1952, 29, 539–543. Kimura, M.; Kimura, K.; Aoki, M.; Shibata, S. Investigation on the Molecular Structures of Titanium Tetrachloride and Zirconium Tetrachloride by Gas Electron Diffraction. Bull. Chem. Soc. Jpn. 1956, 29 (1), 95–100. Avens, L. R.; Cribbs, L. V.; Mills, J. L. Exchange-Reactions of Tetrakis(trifluoromethyl)diphosphine With Pnicogen-Pnicogen, Phosphorus-Hydrogen, and Phosphorus-Chlorine Bonds. Inorg. Chem. 1989, 28 (2), 211–214. Weinhold, F.; Landis, C. R. Valency and Bonding: A Natural Bond Orbital Donor-Acceptor Perspective; Cambridge University Press: Cambridge, 2005. p. ix, 749 p. Bent, H. A. An Appraisal of Valence-Bond Structures and Hybridization in Compounds of the 1st-Row Elements. Chem. Rev. 1961, 61 (3), 275–311. Yamamoto, S.; Nakata, M.; Fukuyama, T.; Kuchitsu, K. Geometrical Structure of Cyclopropane as Studied by Gas Electron-Diffraction and Spectroscopic Data. J. Phys. Chem. 1985, 89 (15), 3298–3302. Reed, Z. D.; Duncan, M. A. Infrared Spectroscopy and Structures of Manganese Carbonyl Cations, Mn(CO)+n (n ¼ 1-9). J. Am. Soc. Mass. Spectrom. 2010, 21 (5), 739–749.

30

Models for Understanding Main Group and Transition Metal Bonding

23. Grushin, V. V.; Marshall, W. J. The Fluoro Analogue of Wilkinson’s Catalyst and Unexpected Ph-Cl Activation. J. Am. Chem. Soc. 2004, 126 (10), 3068–3069. 24. Macgregor, S. A.; Roe, D. C.; Marshall, W. J.; Bloch, K. M.; Bakhmutov, V. I.; Grushin, V. V. The F/Ph Rearrangement Reaction of (Ph3P)3RhF, the Fluoride Congener of Wilkinson’s Catalyst. J. Am. Chem. Soc. 2005, 127 (43), 15304–15321. 25. Gilheany, D. G. No d-Orbitals But Walsh Diagrams and Maybe Banana Bonds – Chemical Bonding in Phosphines, Phosphine Oxides, and Phosphonium Ylides. Chem. Rev. 1994, 94 (5), 1339–1374. 26. Brown, R. K.; Williams, J. M.; Schultz, A. J.; Stucky, G. D.; Ittel, S. D.; Harlow, R. L. Delocalized 2-Electron 3-Center c-h-Metal Interaction – Single-Crystal Neutron (30 and 110 K) and X-ray (298 K) Diffraction Study of Fe(P(OCH3)3)3(ETA 3-C8H13)+ BF−4 . J. Am. Chem. Soc. 1980, 102 (3), 981–987. 27. Harlow, R. L.; McKinney, R. J.; Ittel, S. D. Structure of eta-3-Cyclooctenyltris(trimethyl phosphite)iron(I) – Bonding of the eta-3-Alkenyl Group to 16-Electron, 17-Electron, and 18-Electron ML3 Systems. J. Am. Chem. Soc. 1979, 101 (25), 7496–7504. 28. Wolczanski, P. T. Flipping the Oxidation State Formalism: Charge Distribution in Organometallic Complexes as Reported by Carbon Monoxide. Organometallics 2017, 36 (3), 622–631. 29. Andrei, H. S.; Solca, N.; Dopfer, O. IR Spectrum of the Ethyl Cation: Evidence for the Nonclassical Structure. Angew. Chem. Int. Ed. 2008, 47 (2), 395–397. 30. Raghavachari, K.; Whiteside, R. A.; Pople, J. A.; Schleyer, P. V. Molecular-Orbital Theory of the Electronic-Structure of Organic-Molecules. 40. Structures and Energies of C-1-C-3 Carbocations, Including Effects of Electron Correlation. J. Am. Chem. Soc. 1981, 103 (19), 5649–5657. 31. Trinquier, G. Polymorphism in the Heavier Analogs of the Ethyl Cation. J. Am. Chem. Soc. 1992, 114 (17), 6807–6820. 32. McGrady, G. S.; Turner, J. F. C.; Ibberson, R. M.; Prager, M. Structure of the Trimethylaluminum Dimer as Determined by Powder Neutron Diffraction at Low Temperature. Organometallics 2000, 19 (21), 4398–4401. 33. Noh, S. K.; Heintz, R. A.; Janiak, C.; Sendlinger, S. C.; Theopold, K. H. A Paramagnetic m-Methylene Complex With a Short Cr(III)-Cr(III) Bond. Angew. Chem. Int. Ed. 1990, 29 (7), 775–777. 34. Schmidt, G. F.; Muetterties, E. L.; Beno, M. A.; Williams, J. M. Alkyl-Metal and Aryl-Metal Bond Chemistry in Coordinately Unsaturated Polynuclear Metal-Complexes. Proc. Nat. Acad. Sci. 1981, 78 (3), 1318–1320. 35. Schmidbaur, H.; Tronich, W. Pure Synthesis and Properties of Trialkyl-Alkylidene-Phosphoranes. Chem. Ber. Recl. 1968, 101 (2), 595. 36. Hoffmann, R. Building Bridges Between Inorganic and Organic-Chemistry (Nobel Lecture). Angew. Chem. Int. Ed. 1982, 21 (10), 711–724. 37. Landis, C. R.; Firman, T. K.; Root, D. M.; Cleveland, T. A Valence Bond Perspective on the Molecular Shapes of Simple Metal Alkyls and Hydrides. J. Am. Chem. Soc. 1998, 120 (8), 1842–1854. 38. Firman, T. K.; Landis, C. R. Valence Bond Concepts Applied to the Molecular Mechanics Description of Molecular Shapes. 4. Transition Metals With pi-Bonds. J. Am. Chem. Soc. 2001, 123 (47), 11728–11742. 39. Firman, T. K.; Landis, C. R. Structure and Electron Counting in Ternary Transition Metal Hydrides. J. Am. Chem. Soc. 1998, 120 (48), 12650–12656. 40. Fey, N.; Orpen, A. G.; Harvey, J. N. Building Ligand Knowledge Bases for Organometallic Chemistry: Computational Description of Phosphorus(III)-Donor Ligands and the Metal-Phosphorus Bond. Coord. Chem. Rev. 2009, 253 (5–6), 704–722. 41. Robinson, S. G.; Sigman, M. S. Integrating Electrochemical and Statistical Analysis Tools for Molecular Design and Mechanistic Understanding. Acc. Chem. Res. 2020, 53 (2), 289–299. 42. Santiago, C. B.; Guo, J. Y.; Sigman, M. S. Predictive and Mechanistic Multivariate Linear Regression Models for Reaction Development. Chem. Sci. 2018, 9 (9), 2398–2412. 43. Poe, A. J. Pendent Group Effects, PGEs, in P-Donor Ligands. New J. Chem. 2013, 37 (10), 2957–2964. 44. Fernandez, A.; Reyes, C.; Lee, T. Y.; Prock, A.; Giering, W. P.; Haar, C. M.; Nolan, S. P. Assessing the Stereoelectronic Properties of Pyrrolyl Phosphines and Related Ligands. The Quantitative Analysis of Ligand Effects (QALE). Perkin Trans. 2000, (7); 1349–1357. 45. Fernandez, A. L.; Wilson, M. R.; Prock, A.; Giering, W. P. Evaluation of the Stereoelectronic Parameters of Fluorinated Phosphorus(III) Ligands. The Quantitative Analysis of Ligand Effects (QALE). Organometallics 2001, 20 (16), 3429–3435. 46. Woska, D.; Prock, A.; Giering, W. P. Determination of the Stereoelectronic Parameters of PF3, PCl3, PH3, and P(CH2CH2CN)3. The Quantitative Analysis of Ligand Effects (QALE). Organometallics 2000, 19 (22), 4629–4638. 47. Durand, D. J.; Fey, N. Computational Ligand Descriptors for Catalyst Design. Chem. Rev. 2019, 119 (11), 6561–6594. 48. Evans, L. A.; Fey, N.; Harvey, J. N.; Hose, D.; Lloyd-Jones, G. C.; Murray, P.; Orpen, A. G.; Osborne, R.; Owen-Smith, G. J. J.; Purdie, M. Counterintuitive Kinetics in Tsuji-Trost Allylation: Ion-Pair Partitioning and Implications for Asymmetric Catalysis J. Am. Chem. Soc. 2008, 130 (44), 14471. 49. Jover, J.; Fey, N.; Purdie, M.; Lloyd-Jones, G. C.; Harvey, J. N. A Computational Study of Phosphine Ligand Effects in Suzuki-Miyaura Coupling. J. Mol. Catal. A Chem. 2010, 324 (1–2), 39–47. 50. Tolman, C. A. Steric Effects of Phosphorus Ligands in Organometallic Chemistry and Homogeneous Catalysis. Chem. Rev. 1977, 77 (3), 313–348. 51. Alyea, E. C.; Song, S. Q. Re-examination of the Metal Carbonyl Complex Infrared Parameter, n(co), and Phosphorus Ligand Parameters, pKa, Sigma(chi i) and Sigma Sigma(ph), in Relation to an Evaluation of Sigma and pi Components of M-P Bonds. Comments Inorg. Chem. 1996, 18 (4), 189–221. 52. Bemowski, R. D.; Singh, A. K.; Bajorek, B. J.; DePorre, Y.; Odom, A. L. Effective Donor Abilities of E-t-Bu and EPh (E ¼ 0, S, Se, Te) to a High Valent Transition Metal. Dalton Trans. 2014, 43 (32), 12299–12305. 53. Billow, B. S.; McDaniel, T. J.; Odom, A. L. Quantifying Ligand Effects in High-Oxidation-State Metal Catalysis. Nat. Chem. 2017, 9, 837. 54. DiFranco, S. A.; Maciulis, N. A.; Staples, R. J.; Batrice, R. J.; Odom, A. L. Evaluation of Donor and Steric Properties of Anionic Ligands on High Valent Transition Metals. Inorg. Chem. 2012, 51 (2), 1187–1200.

1.03

Reversible Homolysis of Metal-Carbon Bonds

Maxime Michelas, Christophe Fliedel, and Rinaldo Poli, Laboratoire de Chimie de Coordination, UPR CNRS 8241, Toulouse, France © 2022 Elsevier Ltd. All rights reserved.

1.03.1 Introduction 1.03.2 General aspects of homolytic metal-carbon bond cleavage 1.03.2.1 Energy profile for thermal activation 1.03.2.2 Photoinduced cleavage 1.03.2.3 The “persistent radical effect” 1.03.2.4 Thermal stability 1.03.2.5 Bond cleavage activation parameters 1.03.2.6 Bond formation activation parameters 1.03.2.7 Calorimetric studies of metal-carbon bond strengths 1.03.2.8 Other methods to measure BDFEs/BDEs 1.03.2.8.1 Equilibrium measurements 1.03.2.8.2 Decomposition kinetics 1.03.2.8.3 Electrochemical simulations 1.03.2.9 Computational studies 1.03.3 Reversible metal-carbon bond homolysis in biochemistry 1.03.3.1 Vitamin B12 and derivatives: General aspects 1.03.3.2 Coenzyme B12-dependent enzymatic reactions involving cobalt(III)-carbon bond homolysis 1.03.3.3 Radical S-adenosyl-L-methionine 1.03.3.4 Metal-carbon bond homolysis in other enzymes 1.03.4 Reversible metal-carbon bond homolysis in metal-mediated and -catalyzed organic transformations 1.03.4.1 General aspects of radical reactions in the presence of metals 1.03.4.2 Metal-based radical generations 1.03.4.2.1 By reduction of a polar R–Y bond 1.03.4.2.2 By H atom transfer (HAT) to an alkene 1.03.4.2.3 By metal-carbon bond homolysis 1.03.4.3 Role of metal-carbon bonds in radical reactions 1.03.4.3.1 Hydrogenation 1.03.4.3.2 Dehydrometallation 1.03.4.3.3 Alkyl-hydride reductive elimination 1.03.4.3.4 Dialkyl reductive elimination 1.03.4.3.5 Alkyl transfer to an electrophile 1.03.4.3.6 Oxidation 1.03.4.3.7 Transmetalation 1.03.5 Reversible metal-carbon bond homolysis in controlled radical polymerization 1.03.5.1 General aspects of organometallic-mediated radical polymerization (OMRP) 1.03.5.2 Titanium 1.03.5.3 Vanadium 1.03.5.4 Chromium 1.03.5.5 Molybdenum 1.03.5.6 Manganese and rhenium 1.03.5.7 Iron 1.03.5.8 Ruthenium and osmium 1.03.5.9 Cobalt 1.03.5.9.1 Porphyrin systems 1.03.5.9.2 b-Diketonate systems 1.03.5.9.3 Other planar macrocyclic systems 1.03.5.9.4 Other ligand systems 1.03.5.10 Rhodium 1.03.5.11 Copper Acknowledgment References

31 32 32 33 34 35 36 40 41 41 41 43 43 46 47 47 49 52 53 54 54 56 56 58 59 60 60 61 61 62 62 63 64 64 64 66 68 69 71 72 73 74 74 74 75 76 77 78 78 78 78

Comprehensive Organometallic Chemistry IV

31

https://doi.org/10.1016/B978-0-12-820206-7.00075-5

32

Reversible Homolysis of Metal-Carbon Bonds

1.03.1

Introduction

It is fair to say that metal alkyl and aryl compounds, i.e. compounds where metal and carbon share two electrons in a single bond, are the quintessence of organometallic chemistry. Metal alkyl compounds have been known for a long time, starting from Frankland’s diethylzinc (1848),1 Grignard’s magnesium reagents (1900)2,3 and Gilman’s cuprate reagents in the early 1930s,4 long before organometallic chemistry flourished as a discipline. Many metal alkyl and aryl compounds are quite reactive species, able to engage in a variety of reactions (bond heterolysis, oxidation, migratory insertion, reductive elimination, bond metathesis, etc.), making them delicate compounds to synthesize, isolate and characterize. However, for the same reasons, they are extremely useful and versatile in catalysis. Most of the aforementioned reactions involve 2-electron changes in the metal coordination sphere. This chapter deals with a different process, homolytic bond cleavage, which is a 1-electron process and generates organic radicals. As organic radicals are themselves very reactive species and inevitably disappear through irreversible coupling and disproportionation processes, the metal-carbon bond in organometallic compounds must have sufficient homolytic strength to allow isolation and characterization. Nevertheless, it has become increasingly apparent that homolytically weak metal-carbon bonds, which are prone to reversibly generate transient radical species, are extremely useful to promote unique reactivity. Awareness of the importance of homolytically weak metal-carbon bonds in chemical reactivity is only quite recent and has not been highlighted in any of the previous versions of Comprehensive Organometallic Chemistry. The importance of metal-carbon bond homolysis was initially highlighted in biochemistry, where metal centers in enzymes and cofactors play a crucial role in the physiological regulation of radical reactivity, but has more recently led to growing contributions in the areas of organic chemistry and polymer chemistry. In this article, general aspects related to homolytically weak metal-carbon bonds are first highlighted. Subsequently, separate sections are dedicated to the three above-mentioned areas: biochemistry, organic chemistry and polymer chemistry. Since the more reactive aryl radicals form homolytically stronger bonds with transition metals, homolytic cleavage is mostly relevant for metal alkyl compounds, although all developed concepts also apply in principle to metal-aryl bonds. A generic metal will be abbreviated as Mt in this article, rather than as the more common M, to avoid confusion in the polymerization section where M will indicate a generic monomer. A generic formal oxidation state will be indicated by the symbol x (e.g., Mtx, Mtx+1). Therefore, to avoid confusion, a generic one-electron ligand will be identified by the symbol Y and not by the more commonly used X, but the symbol T will also be used when this species acts as a radical trap. A generic coordination sphere will be identified as L/ (e.g., L/Mtx). This article will make use of the non-SI unit of kcal mol−1 for energy, which is more widely used in thermochemical studies than the SI unit kJ mol−1 (1 kcal ¼ 4.183 kJ).

1.03.2

General aspects of homolytic metal-carbon bond cleavage

1.03.2.1

Energy profile for thermal activation

The reversible homolytic bond cleavage is regulated by the key thermodynamic and kinetic parameters shown in Fig. 1. These parameters may be experimentally assessed and computationally predicted on the enthalpy and/or free energy scales. The Bond Dissociation Enthalpy (BDE) and Bond Dissociation Free Energy (BDFE) define the bond strength from the thermodynamic point of view. The BDE can be experimentally obtained from constant pressure calorimetric measurements, generally requiring the application of thermochemical cycles. The BDFE may be derived through the van’t Hoff relationship from the equilibrium constant K, which may be experimentally obtained, in principle, by equilibrium measurements. Measurements of K at different temperatures

Fig. 1 Energetic profile of metal-carbon bond cleavage.

Reversible Homolysis of Metal-Carbon Bonds

33

also provide an alternative access to the BDE and to the entropy change. Direct measurements of the equilibrium, however, are generally impossible due to the instability of the radical species. Hence, indirect methods are necessary, as will be discussed in Section 1.03.2.8. In terms of kinetics, the activation rate constant (ka) is accessible by measuring the rate of disappearance of the organometallic precursor, provided that the recombination back-reaction (radical deactivation) is removed from the kinetic scheme.5 This may be accomplished by introducing an excess amount of an efficient radical trapping species (T), with which the produced R% species reacts faster than with the metal complex L/Mtx (saturation kinetics) and, at the same time, does not react with the organometallic precursor (Scheme 1). The deactivation rate constant (kda) is not a readily accessible parameter, as its direct measurement requires

Scheme 1 Kinetics approach for the measurement of the activation rate constant.

either monitoring the disappearance of L/Mtx in the presence of a known concentration of radicals (which is possible in flash photolysis or pulse radiolysis experiments), or the measurement of the kda/ktrap competition and the separate knowledge of ktrap. This rate constant can also be estimated from the independent knowledge of ka and K. In the absence of steric impediments, geometrical reorganization, ligand dissociation, or a spin state change, coupling of the organic radical with L/Mtx proceeds at diffusion-limited rates, namely the barrier to the radical deactivation process is very small (estimated as 108 L mol−1 s−1. There seems to be only one report of a variable temperature study, for the addition at low pH of %CH2OH, produced by radiolysis of N2O-saturated aqueous solutions containing CH3OH, to complex [CoII(nta) (H2O)2]− (nta ¼ nitrilotriacetate, N(CH2COO−)3), which yielded DH{ ¼ 4.8  0.5 kcal mol−1 and DS{ ¼ − 4.6  2 cal mol−1 K−1 for measured addition rates of (0.97–4.1)  108 M−1 s−1 in the 7–55  C range.111 The relatively high DH{ value obtained from this study might be associated to the need to displace a water molecule from the coordination sphere. For solvated cations, particularly

Reversible Homolysis of Metal-Carbon Bonds

41

in water, dissociation of the coordinated solvent seems indeed an important step during the Mtx+1-R bond formation process, as suggested by the activation volumes measured by pressure-dependent kinetic studies,112,113 thus suggesting that the diffusion limit hypothesis and use of the viscous flow approximation for DH{da may need reconsideration for systems measured in coordinating solvents. For certain redox-active systems, simulation of the electrochemical data (see Section 1.03.2.8.3) also allows the kda determination with relatively good precision.

1.03.2.7

Calorimetric studies of metal-carbon bond strengths

Thermochemical methods of deriving homolytic Mt-C BDEs have been quite popular. The extraction of Mt-C BDEs from these thermochemical data is however dependent on thermochemical cycles, i.e., on the combination of different measured values resulting in the accumulation of experimental errors, and on a number of approximations and assumptions, such as the effect of solvation in the coordinatively unsaturated complexes. A previous review has outlined the pitfalls of these methods.114 Combustion calorimetry, yielding enthalpies of formation, is complicated for metal-containing substances by, among other factors, the formation of metal oxides having several stoichiometries.115 Several alternative methods, such as reaction calorimetry, heat flux microcalorimetry, differential scanning calorimetry, and photoacoustic calorimetry, have been applied to the study of metal-carbon bond strengths.114,116 In most cases, the process involves the formation or cleavage of more than one bond and thus the extraction of a specific BDE requires the precise knowledge of the others. Solvation effects can be eliminated by moving to gas-phase methods with the help of mass spectrometry, but these methods cannot be easily applied to the analysis of uncharged molecules. A compilation of experimentally determined Mt-C BDEs by calorimetric methods is available in the above-mentioned review.114 However, many reported values are affected by relatively large and sometimes probably underestimated errors. Given this state of affairs and the waning popularity of this line of research, we will not enter into the details of each measurement method. Rather, we illustrate one representative example of how an imprecisely determined BDE has been recently reassessed. The Mn-CF3 BDE in [(CO)5Mn-CF3] was obtained from thermochemical cycles that used microcalorimetric determinations of the enthalpies of sublimation, thermal decomposition, bromination and iodination of the compound, yielding its enthalpy of formation. The combination of this value with the known Mn-Mn BDE in [Mn2(CO)10] (94 kJ mol−1), plus a few assumptions, led to the estimation of the Mn-CF3 BDE as 172  7 kJ mol−1 (41.1  1.7 kcal mol−1).117 A later re-evaluation, based on a new and perceived more precise Mn-Mn BDE in [Mn2(CO)10] (159  21 kJ mol−1), placed the Mn-CF3 BDE at 203  6 kJ mol−1 (48.5  1.4 kcal mol−1).114 An independent photoionization mass spectrometric study that used another thermochemical cycle and other assumptions gave a BDE of 182  11 kJ mol−1 (43.5  2.6 kcal mol−1).114,118 However, all these values are much lower than the kinetically determined activation enthalpy reported in Table 1 (53.8  3.5 kcal mol−1).60 The latter value, which should be rather close to the BDE (difference of 107 M−1 s−1, in agreement with the few values obtained by pulse radiolysis and flash photolysis studies (Section 1.03.2.8.2). When considering only the (presumably more accurate) values obtained in the more recent investigation,130 it appears that secondary radicals are trapped faster than primary ones. The remarkable trapping ability of [CuI(TPMA)]+ for the secondary %CHMeCOOMe radical, (1.3  0.2)  109 107 M−1 s−1 in DMF, was confirmed by additional CV simulations in the presence of different amounts of TEMPO.130 Although so far only applied to alkylcopper(II) system, this method appears of potentially wider applicability. The accurate estimation of the OM equilibrium parameters is limited to systems with labile Mtx+1-R bonds, thus generally not amenable to isolation as pure compounds, at least under standard laboratory conditions. They must be characterized by sufficiently high activation rate constants to have an impact on the CV shape within the timescale of the measurement, but nevertheless strong enough Mtx+1-R bonds to allow the generation of observable amounts of the organometallic species in situ. The range of KOM reported in Table 3 (ca. 10−10–10−6) corresponds to BDFEs in the 8.2–13.6 kcal mol−1 range.

1.03.2.9

Computational studies

Given the general difficulty of obtaining information on the homolytic metal-carbon bond strength from experimental studies, and given the growing performance of computation methods coupled with the increased speed of computers, growing efforts have been devoted to the computational estimation of Mtx+1-R BDEs. Density functional theory (DFT) is the preferred approach in most investigations and has nowadays almost completely supplanted the ab initio methods, because it combines a relatively good accuracy with a lower computational cost. We shall not enter into the technical details of the various available methods and refrain from covering the long and rapidly evolving list of computed Mtx+1-R BDEs. We will limit this overview to a few general considerations, to recommendations for the appropriate use of this approach, and to a discussion of its general utility with a few illustrative examples. A first and most important point is that there is no way to assess the computational error. A computed BDE may be greater or smaller than the true value. Ab initio methods, resting on the variational principle, ensure that the calculated total energy of a system is greater than its true value. Therefore, any improvement of the method can be assessed from the decrease of the computed total energy. However, the BDE of interest is the difference between the energies of L/Mtx+1-R and the sum of L/Mtx and R%. The resulting error may therefore be either positive or negative. Furthermore, the variational principle is not valid for the DFT methods and the error cannot be assessed even for a single system. Hundreds of different functionals have been developed and are available in the toolkits of commercial program packages: local density approximation (LDA) and generalized gradient approximation (GGA) functionals, either pure, hybrid with mixing of a certain percent of exact Hartree-Fock exchange, or spin-component scaled double-hybrid, long-range corrected and/or dispersion-corrected. Very different BDEs may result from the application of different functionals. The effect of the choice of functional is often larger than any correction associated with the physically meaningful solvation effect or with the basis set superposition error (BSSE).132,133 Particular care has to be exercised when calculating complexes than can adopt two or more different electronic configurations. Certain functionals are poorly calibrated for the estimation of the relative stability of spin isomers, overestimating the stability of the higher or lower spin state. A few have been optimized to reproduce the behavior of spin-crossover compounds, e.g., by modulating the percent of exact Hartree-Fock exchange.134 Thus, functionals that adequately describe the L/Mtx+1-R bond homolysis with spin conservation (e.g., from ½ for [L/Mtx+1-R] to 0  ½ for [L/Mtx] + R%), may be unsuitable to describe homolyses with a spin change. For instance, the intensively investigated coenzyme B12, where the diamagnetic [L/CoIII-R] produces two spin doublet fragments, has afforded BDEs varying from 29.5 kcal mol−1 (near the experimentally accepted value) to 15 or more kcal mol−1 lower.135–140 Appropriate consideration of van der Waals interaction (dispersion forces) is crucial. Its neglect may affect the BDE by as much as 12.8 kcal mol−1 for the cobalamin systems138 and such a large effect appears to be general for the Mtx+1-C bond homolysis.137 However, it was also shown that an equally important error, for diamagnetic Mtx+1-R systems like the cobalamins, is the overestimation of the stability of the two fragments by hybrid functionals (e.g., B3LYP), relative to pure functionals (e.g., BP86). This effect was attributed to an overstabilization of the diradical structure from cleavage of the bond.137,141,142 Another striking example

Reversible Homolysis of Metal-Carbon Bonds

47

is provided by the CoIII–C bond in diamagnetic [(acac)2CoIII-CH(OOCCH3)CH3], which yields the CH3(CH3COO)CH% radical (S ¼ ½) and a spin quartet (S ¼ 3/2) [CoII(acac)2] complex. The same comparative study gave quite different BDEs when using different functionals, varying from 9.3 kcal mol−1 (for the hybrid B3LYP functional) to 34.2 kcal mol−1 , (for the dispersion-corrected M06L functional).143 Thus, confidence in the relevance of any computed value must be derived from benchmarking the computational method against any available (and reliable) experimental value, whether this is an equilibrium value (DH , DG ) or a kinetic parameter (e.g., DG{a obtained from ka). Estimation of the entropic contribution to the BDFE also deserves a comment. Entropy is calculated for the isolated molecule in the gas phase from the nuclear Schrödinger equation at the fixed minimum of the electronic potential energy surface, even when the geometry optimization and thermochemical corrections are carried out in the presence of a solvent model (polarizable continuum). It is the sum of vibrational, rotational, translational and spin contributions. Experimental entropies, however, pertain in most cases to a condensed phase where the relatively important translational and rotational modes are partly quenched. This entropy loss is not considered in the standard computational approach. Hence, the computed entropy of a molecular entity is always overestimated with respect to its true value in solution. When a reaction produces an increased number of independent molecules (such as the L/Mtx+1-R bond homolysis), this error is only partly compensated and the overall entropy change is overestimated. Consequently, it is not advisable to compare computed BDFEs with experimental equilibrium data. It is safer to compare computed and experimental BDEs. Because of all the above caveats, the utility of the computational tool for the quantitative assessment of metal-carbon bond strength is rather limited and can never supplant the experimental determination. On the other hand, computational investigations are invaluable in two different types of situations: (i) to rationalize observed (sometimes unexpected) phenomena; (ii) to predict trends when exploring a series of closely related systems. In the first case, the computations provide useful insight and understanding, whereas in the second one they orient the design and experimental development of new systems capable of achieving a desired performance. A notable example, for the first scenario, is an investigation of the BDE in octahedral [L(D)/CoIII-R] systems with an axial donor ligand (D) trans to R.144 After benchmarking against 30 different compounds (yielding a mean deviation of 1.8 kcal mol−1 for the selected functional), systematic modification of the in-plane substituents while keeping R (CH3) and D (pyridine) constant revealed an unexpected negative correlation between the BDE and the spin density at the cobalt atom after the homolysis (bonds are weaker when the free electron is more localized on the Co atom). Analysis of the molecular orbitals and atomic charges revealed that the stabilization effect of the in-plane ligands is larger for the starting material than for the radical, thus rationalizing the unexpected trend. A number of other examples of both kinds of assistance are available in the following sections.

1.03.3

Reversible metal-carbon bond homolysis in biochemistry

The demonstration of reversible metal-carbon bond homolysis in biological systems has long been limited to coenzyme B12-dependent enzymes (Section 1.03.3.2).145,146 However, recent studies have shown that the reaction mechanism of radical S-adenosyl-L-methionine (SAM) enzymes also involves an organometallic intermediate (O), which further generates an active radical via reversible metal-carbon bond homolysis (Section 1.03.3.3).147,148 Although the presence of a metal-carbon bond was established, such as in B12-dependent methyl-transferases in which the methyl transfer occurs by a two-electron pathway,149 or speculated as in certain nickel metalloenzymes,150,151 none of these species undergo a reversible metal-carbon bond homolysis (Section 1.03.3.4).

1.03.3.1

Vitamin B12 and derivatives: General aspects

Vitamin B12 (abbreviated “B12”) and its derivatives, especially those exhibiting a cobalt-carbon bond (Fig. 4), have a crucial role in the human body.152 Therefore, many investigations have been dedicated to their isolation,153 to the study of their biological

Fig. 4 Vitamin B12 derivatives containing a Co–C bond.

48

Reversible Homolysis of Metal-Carbon Bonds

organometallic chemistry154,155 including enzymatic reaction mechanisms,156–159 to the synthesis and reactivity of model compounds including the evaluation of Co-C BDEs,154 and to the investigation of the interactions with other biological substrates.153 The vitamin B12 structure is based on a corrin ring coordinated to the cobalt(III) ion via its four N donors. In the “base-on” mode, the metal is also axially coordinated by a 5,6-dimethylbenzimidazole nucleotide, which is linked to the corrin ring. In certain enzymes, this donor is replaced by the imidazole moiety of an active site histidine residue (“base-off/His-on” mode). In its natural form, the sixth ligand of vitamin B12 is a hydroxo group. However, the most common and widely commercialized form is the air-stable cyano derivative ((CN)Cbl). The structures of cyanocobalamin ((CN)Cbl),160 coenzyme B12 (AdoCbl)161 and methylcobalamin (MeCbl)162 have been confirmed by X-ray crystallography. The total synthesis of (CN)Cbl was achieved in the 1970s by the groups of Woodward and Eschenmoser.163,164 However, a partial synthesis of AdoCbl from a reduced form of B12165 was reported in the early 1960s.166 Different approaches to access B12 derivatives were further developed, which include (i) trapping of an in situ generated carbon-centered radical (generated from a carboxylic acid, VIII and O2) by the reduced form (CoII) of B12 (abbreviated as B12r),167 (ii) reaction between a “super-nucleophilic” CoI or cobalt hydride complex and electrophilic reagents (e.g., alkyl halide),166 or (iii) reduction of aquacobalamin chloride (by formate ions) to a CoII species, followed by alkynyl or ethyl benzyl radical trapping.168,169 The second route is the most commonly used to access organometallic B12-derivatives.154 Co–C bond cleavage and formation are the central steps in the reactions catalyzed by B12 organometallic derivatives in both biological processes (see Section 1.03.3.2) and organic synthesis.170 While both homolytic and heterolytic Co–C bond dissociation modes have been observed in such organometallic reactions, the present article will only focus on the former.171 In the mid-1980s, the groups of Finke and Halpern studied the anaerobic thermolysis of the Co–C bond of AdoCbl in aqueous media or ethylene glycol and evaluated its BDE (see Section 1.03.2.5 and Table 1).82–84 The different studies confirmed that the homolysis of the Co–C bond is the major mode of decomposition at neutral pH and its BDE is approximately 30–35 kcal mol−1. When the thermolysis is performed in the presence of a radical trap (T: TEMPO83 or bis(dimethylglyoximato)cobalt(II)82), both the formation of B12r and the expected Ado-T species could be observed and monitored. In the absence of radical trap, 8,50 anhydro-50 -deoxyadenosine and 50 -deoxyadenosine were observed as the Ado% decomposition products, in addition to B12r.83 Finke and Hay however showed that the decomposition mode is highly sensitive to pH and temperature, while 88% heterolysis was observed at pH 4.0 and 85.0  C, homolysis reached 90% at pH 7.0 and 85.0  C and 97% at pH 7.0 and 110.0  C.84 A recent computational study highlighted that several factors, such as the cage effect (radicals are kept in close proximity), stabilizing van der Waals interactions or geometrical conformations of the protein, are responsible for the great acceleration ( 1012) of the Co–C bond homolysis in the enzymes (glutamate mutase was used as representative example) compared to the isolated coenzyme.172 Later on, Finke and coworkers showed that the thermolysis of complex 3.1, a model complex of the coenzyme B12, led to an equilibrium between the latter and complex 3.2, resulting from the migration of the benzyl group (Scheme 10). Notably, the radical recombination product dibenzyl was not detected, because of the “persistent radical effect”.56 Considerable research has been devoted to understand the promoting effect of the axial ligand on the Co–C bond homolysis in B12-dependent enzymatic reactions, with the goal of rationalizing the switch from the “base-on” to the “base-off/His-on” mode of coenzyme B12 observed in a few enzymes. A comparative study of the effect of phosphine ligands (L) of varying electronic (pKa) and steric (cone angle) properties on the CoIII-R homolysis for compounds [(L)(DH)2Co-CH2Ph] (DH ¼ dimethylglyoxime anion) and [(L)(OEP)Co-CH2Ph] (OEP ¼ octaethylporphyrin dianion) showed that the former family has an inverse linear dependence on the

Scheme 10 Equilibria involved in the cobalt to carbon benzyl migration in a model of coenzyme B12.

Reversible Homolysis of Metal-Carbon Bonds

49

Fig. 5 Dependence of the Co–C bond strength on the electronic (pKa) and steric (cone angle) of the phosphine ligand L in the [(L)(DH)2Co-CH2Ph] and [(L)(OEP)CoCH2Ph] families. Reproduced with permission from Geno, M. K.; Halpern, J. J. Am. Chem. Soc. 1987, 109, 1238–1240. Copyright 1987 American Chemical Society.

steric parameter, whereas the latter has a linear dependence on the electronic parameter (Fig. 5). This reflects the rigidity of the porphyrin ring, where only the ligand basicity affects the bond homolysis through the better stabilization of the higher oxidation state by the stronger electron donor, whereas the flexibility of the glyoximato coordination sphere transmits a dominant steric perturbation, like the flexible corrin ring in the coenzyme B12. The flexibility and electron richness of the supporting ligand may rationalize Nature’s choice of the corrin scaffold, rather than the porphyrin scaffold, for the biological function of the coenzyme.63 Very recently, a series of AdoCbl derivatives with different axial ligand were studied by computational and spectroscopic methods. The results indicate that less basic axial ligands stabilize the reduced form (B12r), in accordance with the “base-off/Hison” mode observed for coenzyme B12-dependent Class I isomerases.173

1.03.3.2

Coenzyme B12-dependent enzymatic reactions involving cobalt(III)-carbon bond homolysis

The salient and common feature of all known coenzyme B12-dependent enzymatic reactions is the Co–C bond homolytic cleavage that affords the 50 -deoxy-50 -adenosyl (Ado%) reactive radical and B12r.174 Apart from ribonucleotide reductases (see below), the general mechanism of these reactions rests on a 1,2-isomerization of protein-bound radicals, as depicted in Scheme 11.175 The catalytic cycle begins with a first hydrogen abstraction on the substrate by the initially released Ado% radical, leading to Ado-H and a

Scheme 11 General catalytic cycle of coenzyme B12–dependent enzymatic reactions. Ado% ¼ 50 -deoxy-50 -adenosyl; B12r ¼ reduced (CoII) cobalamin. The X group (box) and the bond (bold) involved in the 1,2-isomerization reaction are highlighted for clarity.

50

Reversible Homolysis of Metal-Carbon Bonds

new radical species (substrate radical). The latter then undergoes a 1,2-rearrangement of the X group and Ado-H subsequently transfers the hydrogen atom to the product radical to afford the isomerization product and regenerate AdoCbl. This typical reactivity brought Halpern to consider AdoCbl as a “reversible free (carbon-centered) radical carrier.”122 An important point to mention is that the energetics of Ado% formation are unfavorable, as the Co-C BDE was estimated to be 34.5  1.8 kcal mol−1 for the base-off structure.176 Therefore, interactions with the enzyme are crucial for overcoming this barrier. Indeed, as a consequence to substrate binding, AdoCbl readily undergoes homolytic cleavage and substrate-based radicals accumulate on the enzyme during turnover, which indicate that the equilibrium constant for homolysis is now close to unity.177,178 To date, 10 enzymatic reactions requiring the action of the coenzyme B12 (AdoCbl) as cofactor have been identified and are classified into three groups.156,179 The first series of reactions are carbon skeleton rearrangements, Scheme 12, which consist of the exchange of two vicinal groups (one H and one organic moiety) in a (pseudo)intramolecular fashion. In these carbon skeleton mutases (also called Class I isomerases), the B12-cofactor is bound in the “base-off/His-on” mode (see Section 1.03.3.1). Glutamate mutase, the first enzyme for which the action of AdoCbl was shown, realizes the conversion of glutamate to b-methyl aspartate (entry 1 in Scheme 12).180 The cobalt-carbon bond homolysis is accelerated in the enzyme by a factor of  1012, as shown by the rate of B12r formation upon mixing AdoCbl and the substrate or product (reversible reaction).90 In addition, a large kinetic isotope effect (KIE) supports a mechanism in which the Co–C bond homolysis is kinetically coupled to the substrate H (or D) abstraction. In the same series of reactions, methylmalonyl-CoA mutase (MCM) interconverts R-methylmalonyl-CoA and succinyl-CoA (entry 2 in Scheme 12). An examination of the kinetic and thermodynamic parameters associated with the Co–C bond homolysis revealed an enhancement of the homolysis rate for the enzyme-bound vs. free cofactor,178 with a transition state barrier lowering by 17 kcal mol−1. As for the glutamate mutase, a large KIE on the B12r formation was observed when using a D-labeled substrate, suggesting that the Co–C bond homolysis and H atom abstraction are coupled.91,92 Computational studies supported the generation of a “free” Ado% radical,181 or gave comparable activation barriers for the stepwise (Ado% radical formation followed by H-abstraction) and coupled pathways leading to the formation of B12r,182 but only the coupled pathway can rationalize the observed KIE. By studying the mechanism of inhibition of MCM with itaconyl-CoA, Banerjee and coworkers were able to characterize a diradical species, which is composed of a tertiary carbon-centered radical 6 A˚ away from a cobalt-centered one (B12r), by EPR and X-ray crystallography.183 This species results from the trapping of the Ado% radical by the inhibitor. Therefore, this study provides insights into how MCM controls the radical trajectory during catalysis to promote the desired reaction and suppress side reactions. Two other coenzyme B12-dependent carbon skeleton mutases are known: (1) the a-methylene glutarate mutase that catalyzes the reversible rearrangement of 2-methylene-glutarate and (R)-3-methylitaconate and the isobutyryl-CoA mutase that catalyzes the interconversion of isobutyryl-CoA and n-butyryl-CoA (entries 3 and 4, respectively, in Scheme 12).155

Scheme 12 Skeletal enzymatic rearrangements catalyzed by AdoCbl. By analogy with Scheme 11, the X group (box) and the bond (bold) involved in the 1,2-rearrangement reactions are highlighted for clarity.

AdoCbl-dependent enzymes also catalyze elimination reactions of water and ammonia (entries 5 to 7 in Scheme 13).157,184 These Class II eliminases bind the B12-cofactor in the “base-on” mode (see Section 1.03.3.1). Diol dehydratase (DD) catalyzes the dehydration of ethylene glycol (ethane-1,2-diol) and propylene glycol (propane-1,2-diol, entry 5 in Scheme 13) to form acetaldehyde and propionaldehyde, respectively. In the same way, glycerol dehydratase converts glycerol (propane-1,2,3-triol) to 3-hydroxypropanal (entry 6 in Scheme 13). The proposed reaction mechanisms follow the one depicted in Scheme 11, with elimination of a water molecule as an additional step.185,186 In the 1970s, Abeles and coworkers already evidenced, by EPR, the formation of carbon-based radicals by reaction between propane-1,2-diol and diol dehydratase.187 In a recent EPR study, the 1,2-propanediol-1-yl radical, resulting from the H-abstraction of propane-1,2-diol by the Ado% radical (1st step after AdoCbl

Reversible Homolysis of Metal-Carbon Bonds

51

Scheme 13 Enzymatic tandem rearrangement and elimination reactions catalyzed by AdoCbl. By analogy with Scheme 11, the X group (box) and the bond (bold) involved in the 1,2-rearrangement reactions are highlighted for clarity.

homolysis), was identified as an intermediate in the catalytic cycle.188 Computational studies also revealed that the active-site amino acid residues of diol dehydratase influence the catalytic process. Indeed, H-bonds accelerate the cleavage of the Co–C bond and control access of Ado% to the pro-S hydrogen atom of propane-1,2-diol.189 A recent study highlighted the crucial role of diol dehydratase for the reactivation of the damaged cofactors, i.e., when AdoCbl is not regenerated at the end of the catalytic cycle, in glycerol- and O2-inactivated holoenzymes.190 Ethanolamine ammonia lyase (EAL) catalyzes the conversion of ethanolamine (2-hydroxyethylamine) to acetaldehyde, with the loss of one molecule of ammonia (entry 7 in Scheme 13). EPR studies on model reactions catalyzed by DD191 and EAL,192 using an AdoCbl with axially bonded 15N-labeled 5,6-dimethylbenzimidazole, unambiguously showed that both enzymes bind B12 in its “base-on” form. Recent kinetic studies allowed identifying the intermediates and drawing the sequence of the EAL catalytic cycle as follow: (i) radical pair separation; (ii) hydrogen atom transfer from the a-OH methylene group of ethanolamine to Ado%; (iii) 1,2-rearrangement of the NH3 group to the site of the radical; (iv) hydrogen transfer from Ado-H; (v) radical pair recombination; and (vi) release of products (NH3/NH+4 and acetaldehyde) concomitant with substrate binding.193,194 This result is in line with previous computational studies.195–198 Moreover, the latter studies, together with the crystal structures of EAL complexed with B12 analogs and substrates and recent spectroscopic investigations of the EAL-mediated catalytic cycle suggest that the migrating group is almost fully protonated (R-NH+3) due to interactions between the substrate and active-sites residues.199,200 The third type of reactions catalyzed by AdoCbl-dependent enzymes are amino acid rearrangements, which consist in the migration of o-amino groups to the adjacent carbon atom (entries 8–10 in Scheme 14).159,184 The B12-cofactor is bound in the “base-off/His-on” mode in aminomutases (see Section 1.03.3.1). To date, two B12-dependent aminomutases are known: D-ornithine 4,5-aminomutase (OAM)201 and lysine 5,6-aminomutase (LAM).202 OAM catalyzes the D-ornithine conversion to (2R,4S)-diaminovaleric acid (entry 8 in Scheme 14)177 and LAM converts a- and (3S)-b-lysine to 2,5-diaminohexanoic acid and (3S,5S)-diaminohexanoic acid, respectively (entries 9 and 10, respectively in Scheme 14).203 Both enzymes require the cooperation of pyridoxal 50 -phosphate (PLP or vitamin B6) as an additional co-factor, which binds the substrate by substitution of the internal lysine to form a Schiff base.204 Substrate binding initiates a large domain motion, from an open to a closed and catalytically active structure, which is required to bring AdoCbl, PLP and the substrate into close proximity, and to allow the 1,2-amino group migration.205,206 The catalytic cycle then proceeds as follows. After hydrogen atom abstraction by Ado%, the radical substrate isomerization occurs via an azacyclopropylcarbinyl radical,207,208 and then Ado-H delivers the H atom to the product radical, the product is released, Ado%/AdoCbl is regenerated and PLP binds again the residual lysine group.

Scheme 14 Enzymatic 1,2-rearrangements of amino acids catalyzed by AdoCbl. By analogy with Scheme 11, the X group (box) and the bond (bold) involved in the 1,2-rearrangement reactions are highlighted for clarity.

52

Reversible Homolysis of Metal-Carbon Bonds

Ribonucleotide reductases (RNR) catalyze the transformation of ribonucleotides to their deoxyribonucleotide analogs and are, therefore, essential for DNA synthesis. These enzymes are categorized in three classes, noted Class I to Class III, which all catalyze the same reaction but differ by the nature of the implicated cofactor. Class II RNR are B12-dependent enzymes and involve reversible Co–C bond homolysis. Coenzyme B12 is bound in the “base-on” mode in these enzymes. Several comprehensive reviews dedicated to RNR (different classes, structures, mechanisms, etc.) are available.209–211 The proposed catalytic cycle for B12-dependent ribonucleotide reductases (Class II RNR) is depicted in Scheme 15 and is composed of the following steps.212 Initially, homolytic cleavage of the Co–C bond of AdoCbl affords the Ado% radical that abstracts the S-H hydrogen atom of a cysteine residue (Cys408) to afford a thiyl radical.213 EPR studies on a Class II RNR, containing a specifically deuterated Cys408 residue, provided evidence for close proximity between the thiyl radical and the CoII center of the resulting B12r.214,215 This key structural feature was further confirmed by X-ray crystallography.216,217 The thiyl radical then abstract the 30 -hydrogen atom from the ribose ring of the substrate, generating a substrate radical,218 which is then converted into the corresponding 20 -ketyl radical via the loss of a water molecule.219 Subsequently, two cysteine residues, different from Cys408, are oxidized to a disulfide (cystine) while transferring two hydrogen atoms to the substrate. The radical product abstracts then a hydrogen atom from Cys408, affording the product and regenerating the corresponding thiyl radical, which in turn abstracts an H from Ado-H. The latter step affords again the Ado% radical that is trapped by the CoII center to reform AdoCbl after each turnover.212

Scheme 15 Catalytic cycle of the reaction catalyzed by coenzyme B12-dependent ribonucleotide reductases. Pi-Pi-Pi ¼ P3O10.

1.03.3.3

Radical S-adenosyl-L-methionine

Radical S-adenosyl-L-methionine (SAM) enzymes, a “superfamily” comprising several thousand species, possess a [4Fe-4S]+ cluster promoting reductive cleavage and release of the Ado● radical, which initiates over 80 different radical reactions.148,220 In a particularly informative study for pyruvate formate-lyase activating enzyme (PFL-AE/SAM), it was possible to trap a catalytically relevant intermediate, named O, by rapid freeze-quenching (RFQ).221 EPR and 13C, 57Fe electron nuclear double-resonance (ENDOR) spectroscopic techniques established that O has trapped the Ado● radical, resulting from a C–S bond cleavage in SAM, through formation of a Fe–C bond to the unique iron center (the one not coordinated by cysteine) of the [4Fe-4S] cluster with the Ado 50 -carbon, as depicted in Scheme 16. The mechanism of formation of O, which is not yet clearly established, may follow one of the following paths: (i) a two-step path involving SAM reductive cleavage followed by Ado● trapping by the unique Fe center of the [4Fe-4S]2+ cluster, or a single step via either (ii) direct nucleophilic attack of the unique Fe center of the [4Fe-4S]2+ cluster at the SAM 50 -C atom or (iii) concerted reductive cleavage/O formation initiated by interaction between the Fe and the S atoms.222 However, the orientation of the S-C(50 ) bond in the active site of the radical SAM enzymes is not conducive to Fe attack, making path (ii) less likely. Formation of O is not limited to PFL-AE/SAM: later, O was also identified as an intermediate during reactions catalyzed by several other radical SAM enzymes.222 In addition, homolytic S-C50 bond cleavage promoted by photo-induced electron transfer at 12 K has unequivocally revealed the nature of the elusive Ado● radical through a combined EPR/ENDOR and DFT investigation.223 Following this finding, synthetic models of O have been reported and studied, demonstrating reversible Fe–C bond homolysis in protein-free systems.224,225

Reversible Homolysis of Metal-Carbon Bonds

53

Scheme 16 Activation mechanism of radical SAM enzymes involving the formation of the organometallic intermediate O, which undergoes Fe-C homolysis with Ado● release.

Using the RFQ method and EPR and ENDOR analytical techniques, another “O–like” organometallic species was recently identified as intermediate in the diphthamide biosynthesis (3.3 in Scheme 17).226 This 3-amino-3-carboxypropyl-[4Fe-4S]3+ organometallic species may be formed stepwise, via homolytic reductive cleavage of SAM, trapping of the 3-amino-3-carboxypropyl (ACP) radical by the unique Fe of the [4Fe-4S]+ cluster and release of the 50 -methylthioadenosine (MTA) molecule, or in a concerted one-step-two-electron transfer by nucleophilic attack of the Fe on the SAM Cg-Met atom. In this case, structural parameters implicate the second pathway. This intermediate then provides the ACP● radical (3.4) via Fe–C bond homolysis, leading to the diphthamide synthesis (Scheme 17).

Scheme 17 Proposed mechanism for the two first steps of the diphthamide biosynthesis performed by radical SAM enzymes, highlighting the formation of the organometallic intermediate I and its reactivity via Fe–C bond homolysis. MTA ¼ 50 -methylthioadenosine, EF2 ¼ eukaryotic elongation factor 2 (protein).

1.03.3.4

Metal-carbon bond homolysis in other enzymes

Acetyl-CoA synthase is a nickel-containing enzyme that forms, together with carbon monoxide dehydrogenase, the bifunctional acetyl-CoA synthase/carbon monoxide dehydrogenase (ACS/CODH) enzyme. This bifunctional enzyme catalyzes the synthesis of acetyl-coenzyme A (Acetyl-CoA or H3C(CO)CoA) via the Wood-Ljungdahl pathway (Eq. 1, Scheme 18).151 The reaction involves coenzyme A (CoA), CO and the transfer of a methyl group from MeCbl (Fig. 4), the latter being part of the corrinoid iron-sulfur protein (Co-FeSP). Although the formation and transfer of a methyl radical from cobalt to nickel was found plausible according to model studies,227,228 the authors themselves admitted that this reactivity could not occur in the biological system because of the low reduction potential of the CoIII/CoII couple (below −1 V).229,230

54

Reversible Homolysis of Metal-Carbon Bonds

Scheme 18 Global reactions of the Acetyl-CoA (Eq. 1) and methane (Eq. 2) synthesis, mediated by nickel-containing enzymes.

Methyl-coenzyme M reductase (MCR), which contains a square-planar nickel corphin cofactor (F430), catalyzes the reversible transformation of methyl thioether methyl-coenzyme M (H3C-S-CoM) and thiol coenzyme B (CoB-SH) to the corresponding heterodisulfide (CoB-S-S-CoM) and methane (Eq. 2, Scheme 18).145 A mechanism involving transfer of a methyl radical, generated from an activated CoB-SS%(CH3)-CoM species, to nickel(I), to form a methylnickel(II) complex was proposed in the 90s.231,232 However, although the reaction mechanism is not yet fully elucidated, the nickel-methyl intermediate does not react at a kinetically competent rate, and is unlikely to be an intermediate in catalysis.233

1.03.4

Reversible metal-carbon bond homolysis in metal-mediated and -catalyzed organic transformations

Given the extensive and rapidly growing literature, this section cannot provide a comprehensive coverage. Rather, it will highlight a few general principles of radical reactivity in the presence of metals, show how radicals can be produced by metal complexes and survey the various types of processes encountered in metal-mediated or catalyzed organic synthesis via radicals, highlighting, where necessary, possible misconceptions and mechanistic booby traps.

1.03.4.1

General aspects of radical reactions in the presence of metals

In organic synthesis, reactions involving radicals offer interesting functional group tolerance surpassing those associated with nucleophilic and electrophilic reagents and therefore attract growing attention.234 Metal complexes occupy an important role in this area, especially when used in substoichiometric amounts. The most useful complexes are those of 3d metals,235 which can promote one-electron processes more easily than their heavier congeners. Fe236–241 and Co239–243 are most intensively investigated, although other 3d metals (Ti,244–246 V, Cr,247 Mn240,248–251 and Cu252) and certain heavier metals are also of interest. Many radical reactions follow a chain mechanism. The substrate is converted into a radical either by interaction with a primary radical generated from an initiator, or directly, by one of a number of different stimuli (e.g., thermal, photochemical, redox). The substrate radical then transforms into the product radical, which is turned into the stable product by exchange with the substrate (chain transfer) to produce a new substrate radical, Scheme 19A. The chain efficiency depends on the rate of the propagation steps relative to the initiation and termination rates. A metal complex may simply promote initiation in a stoichiometric or catalytic fashion. In that case, the radical reaction is not metal-catalyzed. Rather, it is accelerated through an increased initiation efficiency (greater steady-state radical concentration). This effect can be defined as “smart initiation.”253 A metal complex may also lower the activation barrier of the propagating steps (a true catalytic effect) and/or alter the reaction selectivity through specific interactions with the radical intermediates. These effects generally result from donor-acceptor interactions of the radical and the metal (Lewis acid/base catalysis), not through formation of a metal-carbon s bond. Finally, the metal complex may lower the steady-state radical concentration, hence reduce the impact of bimolecular terminations and improve the chain efficiency, through the installment of a moderating equilibrium with a dormant species (the “persistent radical effect”) and can do so by either atom transfer or direct bond

Scheme 19 Formation of metal-carbon bonds in chain (A) and non-chain (B) radical reactions.

Reversible Homolysis of Metal-Carbon Bonds

55

formation. In this article, we are only concerned with the moderating role of a metal complex through reversible metal-carbon bond formation. A metal complex may also allow non-chain radical reactions, which would normally not take place, by combining a “smart initiation” with stoichiometric radical transformations and a productive termination. The substrate radical is generated from the stable substrate by interaction with the metal complex in a one-electron process where the metal is formally reduced or oxidized. The new complex is then able to efficiently trap the product radical and convert it to the stable product, ideally regenerating the initial metal complex and thus closing a catalytic cycle, Scheme 19B. A spontaneous radical chain reaction with “smart initiation” turns into a metal-catalyzed non-chain process if the product radical quenching by the metal complex is more efficient than chain transfer to the substrate. Mechanistic proposals, particularly in terms of the role of the metal complex, are abundant but seldom accompanied by careful mechanistic investigations. Specifically, with the notable exception of cobalt-catalyzed reaction (vide infra), the implication of organometallic intermediates and of metal-carbon bond homolysis is not often invoked and even less frequently demonstrated. This is understandable, given that compounds with homolytically weak bonds may not accumulate in sufficient amounts for detection. In the absence of detailed experimental or computational (ideally both) mechanistic investigations, care should be exercised when proposing a reaction pathway, particularly when the chemical composition and/or oxidation state of a proposed intermediates is undocumented. Not only has reversible metal-carbon bond formation seldom been considered as a useful moderating phenomenon, it has even been sometimes portrayed as an “unwanted” side reaction, competing with the desired “productive” radical reaction. This overlooks the potential beneficial role of reversible metal-carbon bond formation for limiting the undesired self-termination events. The reaction rates of productive radical transformations scale with the L/Mtx+1-R BDE and inversely with the radical stabilization. At one extreme of the reactivity scale, highly stabilized (unreactive) radicals do not bind to metals, hence no moderating effect is present, but self-terminations also have low impact. This is what happens, for instance, in the [Co2(CO)8]-catalyzed hydrogenation of anthracene and derivatives 4.1, Scheme 20A, where the radical produced by the first hydrogen atom transfer (HAT) reaction is presumably not trapped by [Co(CO)4]% and proceeds to abstract an H atom from a second [HCo(CO)4] molecule to yield the 9,10-dihydroanthracene 4.2 with high selectivity.254 A related example is the stoichiometric hydrogenation of a-methylstyrene by [HMn(CO)5], where the PhC%(CH3)2 radical produced by the initial HAT reaction has no affinity for [Mn(CO)5]%. However, styrene yields small equilibrium amounts of the organometallic coupling product, [(CO)5Mn-CHMePh].255 At the opposite extreme are highly reactive radicals, e.g., non-stabilized alkyls, where moderation is crucial. An illustrative example is provided by the isomerization of 5-hexenyl cobalt(III) complexes, 4.3 in Scheme 20B, to cyclopentylmethyl isomers 4.4.256 The reactants and products have rather strong CoIII–R bonds, hence they are isolable and thermally stable. An initial stimulus produces the 5-hexenyl radical, which spontaneously cyclizes and the cyclopentylmethyl radical is trapped to yield the isomerized product. The reactant can also release its 5-hexenyl radical by exchange with the cyclized radical product (chain process). The presence of self-terminations was not mentioned, but the mass balance of the two organocobalt(III) isomers was always >80%. The moderating role of L/CoII is therefore quite evident. For intermediate situations of increasing radical reactivity and associated metal-carbon bond strengths, metal-carbon bond formation plays a growing role in controlling the selectivity of desired radical reaction versus the undesired selfterminations.

Scheme 20 Two examples of radical reactions with very different radical reactivity: (A) stabilized radical; (B) reactive radical.

56

Reversible Homolysis of Metal-Carbon Bonds

1.03.4.2

Metal-based radical generations

In this section, we analyze how radical formation may be induced in the presence of metal complexes and the interplay, for each of them, with metal-carbon bond formation.

1.03.4.2.1

By reduction of a polar R–Y bond

Substrates containing a polar R–Y bond (typically Y ¼ halide) can be activated by reducing metal complexes able to promote one-electron processes. This process may occur in two distinct ways. 1.03.4.2.1.1 By atom/group transfer (AT/GT) L/Mtx complexes able to increase the metal oxidation state and coordination number by one unit may abstract the Y atom to generate R% and L/Mtx+1-Y (Scheme 21, a). A well-known example is the metal-catalyzed Kharasch addition to alkenes (also called atom transfer radical addition, ATRA),257 which has been extended to atom transfer radical polymerization (ATRP). When radicals are generated by this method, the formation of organometallic L/Mtx+1-R intermediates or dormant species (Scheme 21, b) requires radical diffusion away from the {R%,L/Mtx+1-Y} cage followed by encounter with a second L/Mtx molecule. Reverse AT trapping by L/Mtx+1-Y occurs directly within the radical cage and is a more likely moderating mechanism. However, organometallic species were shown to form in certain ATRP processes (see Section 1.03.5) when the L/Mtx+1-R bond strength is significant. The combination of AT activation and radical trapping by metal is 1-electron oxidative addition (Scheme 21, c), which is well-documented in organometallic chemistry.258 Therefore, the possible formation of L/Mtx+1-R after AT from R-Y + L/Mtx should not be neglected. In addition, the radical may also be trapped directly by the L/Mtx+1-Y product (radical rebound), if the metal can further increase its oxidation state (Scheme 21, d). In this case, the obtained product is equivalent to that of a direct 2-electron oxidative addition

Scheme 21 Activation of a polar R–Y bond by atom transfer.

(Scheme 21, e). In addition to halogen atoms, certain pseudo-halogens, such as carbamates (S2CNR2) and thiocyanate, can also be transferred to L/Mtx.259 Oxiranes 4.5 are particular R-Y substrates, where bond cleavage produces a radical 4.6 which remains chemically linked to the oxidized metal complex, Scheme 22. The radical epoxide ring opening was first developed with [Cp2TiIIICl] and has led to a host of useful transformations,260 but may also be promoted by other systems such as [FeIICl2(dppe)2]/Zn.261,262 A possible moderation of the subsequent radical processes through reversible bond formation with a second L/Mtx molecule to yield 4.7 is not often invoked, although the formation of a 1:2 epoxide/titanium adduct, in the absence of other reagents, was described in the original study260 and also plays a key role in controlled radical polymerization with [Cp2TiCl]/oxirane initiation (see Section 1.03.5.2).

Scheme 22 Activation of oxiranes.

1.03.4.2.1.2 By electron transfer (ET) If the metal complex is strongly reducing, it may transfer an electron to R-Y by outer-sphere electron transfer (OSET) to generate an intermediate radical anion [R-Y]-•, which then eliminates Y− by rapid bond homolysis, Scheme 23. The liberated Y− may recombine with the oxidized complex, especially if this is positively charged ([L/Mtx+1](n+1)+ with n ¼ 0), to yield the same product of AT activation. For negatively charged activating complexes (n ¼ − 1), yielding neutral (L/Mtx+1), this recombination is not expected in principle.

Reversible Homolysis of Metal-Carbon Bonds

57

Scheme 23 Activation of a polar R–Y bond by electron transfer.

As in the AT activation, formation of an organometallic dormant species ([L/Mtx+1-R]n+) requires radical diffusion away from the % {R ,[L/Mtx+1](n+1)+} or {R%,[L/Mtx+1-Y]n+} cage and encounter with a new [L/Mtx]n+ molecule. An organometallic dormant species may form efficiently if the radical rebounds and forms an (x + 2) oxidation state metal complex. This is what happens, for instance, in many processes mediated or catalyzed by an anionic L/CoI species (often generated in situ by reduction of a neutral L/CoII precursor), where L ¼ porphyrin, corrin, tetradentate Schiff (e.g., salen or salophen), or bis(dimethylglyoximato).242,263–273 The ET step produces the radical, Y− and L/CoII and the radical rebound yields a relatively stable L/CoIII-R species. Trapping of primary and/or secondary radicals by L/CoII has been proposed in several catalytic studies with thermal activation. In other cases, the CoIII–C bond is too strong to allow catalytic turnover under thermal conditions, but bond homolysis occurs in the presence of light.272 Note that the “supernucleophilic” anionic L/CoI complexes may also promote SN2 reactions, directly yielding L/CoIII-R in one step.274 The relative aptitude toward SN2 vs. OSET/rebound depends on the nature of L, Y and R substituents. ET occurs more favorably for the more easily reducible iodides and for more substituted R (secondary > primary).275 An early study, for instance, showed that cobaloxime(I) reacts with 2-allyloxyethyl derivatives 4.8 (Scheme 24) to yield a mixture of the two expected non-cyclized (4.9) and cyclized (4.10) products in different proportions depending on the nature of the Y group and R substituents.264

Scheme 24 Reaction of the cobaloxime(I) anion with various allyloxyethyl derivatives.264

In another early example, the formally Fe0 center in [Cp(dppe)Fe-MgBr] 4.11 (Scheme 25) reacts with 1-bromo-5-hexene to yield [Cp(dppe)FeII-CH2-c-C5H9] (4.12) by OSET/cyclization/rebound, whereas the reaction between [Cp(dppe)Fe-Br] and 5-hexen-1-yl Grignard produces the expected SN2 product [Cp(dppe)FeII-(CH2)4CH]CH2] (4.13).276 For FeI or FeII-promoted radical cyclizations, formation of FeII–C or FeIII–C bonds has occasionally been invoked without direct evidence or computational support.

58

Reversible Homolysis of Metal-Carbon Bonds

Scheme 25 OSET/cyclization vs. SN2 for a half-sandwich Fe system.

Certain metal complexes were shown to activate R-Y even though they do not have sufficient reducing power to do so, as demonstrated electrochemically. Therefore, these activations may follow a different pathway from OSET. An inner sphere ET pathway (ISET) may be facilitated by a substrate-catalyst interaction. For instance, [FeIHCl(dppe)2]− and [(Z1-H3BH)FeICl (MeCN)4]−, generated respectively from [FeIICl2(dppe)2] or FeCl2/MeCN and NaBH4, are able to activate allyloxy-substituted organic halides 4.14 (Scheme 26) and promote radical cyclizations related to those of Scheme 24.277,278 For unsaturated substrates, coordination of the unsaturated function as in 4.15 seems possible, since highly reduced metals are p-basic. It is then possible to conceive spin delocalization onto the non-innocent alkene ligand,279 followed by radical reactivity before final reduction to the observed product 4.16.

Scheme 26 Possible radical cyclization pathways for insufficiently reducing FeI metal complexes.

1.03.4.2.2

By H atom transfer (HAT) to an alkene

Unsaturated substrates may be converted to radicals by HAT using a metal hydride complex L/Mtx+1-H, Scheme 27. The latter is often a very reactive species generated in situ from a stable L/Mtx+1-Y precursor (halide, alkoxide, acetylacetonate, etc.) and a main group hydride (silane, stannane, borane, etc.).240,280 An early report of this activation method has already been described in Section 1.03.4.1 (Scheme 20A).254 This radical generation method produces a formally reduced metal complex, L/Mtx. Thus, contrary to the polar R–Y bond activation examined in the previous section, the produced metal species has a suitable electronic configuration to directly bind the radical within the HAT products radical cage 4.17 (Scheme 27) to yield the organometallic species 4.18, if the Mtx+1-C BDE is sufficient. The alternative olefin coordination to yield 4.19 followed by migratory insertion is a well-known 2-electron reaction sequence in many catalytic processes, such as olefin hydrogenation and hydroelementations. The HAT pathway does not require the availability of a vacant coordination site and is preferred when the L/Mtx+1-H and L/Mtx+1-R bonds are homolytically weak.281

Reversible Homolysis of Metal-Carbon Bonds

59

Scheme 27 Activation of an electron-rich alkene by hydrogen atom transfer (HAT).

The L/Mtx+1-H ability as H atom donor is directly related to the Mt–H bond strength, on which much information has been gathered.114,282,283 The lower the BDE, the greater the range of alkenes that can be activated. An efficient system, for instance, is Fe(acac)3/PhSiH3,284,285 where the Fe-H BDE in the [(acac)2FeIII-H] active species was estimated as 17 kcal mol−1 using computations (with the caveats described above).286

1.03.4.2.3

By metal-carbon bond homolysis

There is of course a third and simpler way to generate radicals, namely by direct homolysis of metal-carbon bonds in stable organometallic compounds. If the generated radical undergoes a chemical change, a net stoichiometric reaction results. This pathway, which is reminiscent of that for B12 systems discussed in Section 1.03.3, corresponds to the bottom part of Scheme 19B and is illustrated more generally in Scheme 28.

Scheme 28 Stoichiometric isomerization of organometallic compounds.

This chemical transformation is not of interest per se in organic synthesis, but represents a model for steps that may occur within metal-catalyzed radical reactions. One such process has already been highlighted in Section 1.03.4.1, Scheme 20B.256 The importance of a moderating effect via CoIII–R bond formation was also demonstrated for the alkyl-alkenyl radical cross-coupling by cobaloxime activation in the presence of a-ethoxyacrylonitrile (Scheme 29).287 In the absence of alkene, the photolysis of alkylcobaloxime 4.20 produces the unsaturated product RCH]CH2 and H2 selectively, with only little of the expected self-termination products, because the radical concentration is kept low by the reversible CoIII–C bond formation. In the presence of alkene, both RCH]CH2 and the cross-coupling product 4.21 are obtained, but the former increases at longer reaction times because of the [CoII(dmgH)2py] build-up by occasional terminations. Conversely, the addition of [CoII(dmgH)2py] scavengers tips

Scheme 29 Moderating effect in cross-coupling from an alkylcobaloxime complex.

60

Reversible Homolysis of Metal-Carbon Bonds

the balance toward cross-coupling. When the [CoII(dmgH)2py] concentration is too low, however, polymerization is favored. Other investigations from alkyl-269,288–295 and acyl-296–299 cobalt(III) complexes (either isolated or generated in situ) have served to confirm the involvement of organocobalt(III) species in several cobalt-catalyzed radical processes.

1.03.4.3

Role of metal-carbon bonds in radical reactions

Whether metal-carbon bond formation occurs in radical chain processes or in metal-catalyzed non-chain reactions, it may simply introduce a moderating effect if the dormant organometallic species is unable to undergo other chemical transformations. However, the latter may also further react, thus playing a direct role in metal-mediated radical reactions. If the follow-up reaction is faster than bond homolysis, the latter is suppressed. However, certain follow-up processes may also be triggered by bond homolysis and a few misconceptions in this regard exist in the literature.

1.03.4.3.1

Hydrogenation

Alkene hydrogenation is typically catalyzed by transition metal complexes by cycles involving only 2-electron processes, but certain systems for which the L/Mtx+1-H and L/Mtx+1-R bonds are homolytically weak have been shown to promote radical pathways. The initial olefin insertion may occur by HAT/rebound, as suggested for the [CoII(CN)5]3−-catalyzed hydrogenation of certain alkenes,300,301 whereas no rebound apparently occurs when more stabilized HAT-generated radicals are formed together with [Co0(CO)4] or [Mn0(CO)5].254,255 The intermediate radical, possibly in equilibrium with Mtx+1-R, is then most likely converted to the hydrogenated product by HAT from a second L/Mtx+1-H molecule (e.g., see Scheme 20A) and the released L/Mtx regenerates L/Mtx+1-H by reaction with H2. This radical mechanism may be less common than the typical coordination/migration mechanism of hydrogenation, but certain hydrogenation catalysts, particularly those involving 3d metals and weak L/Mtx+1-R bonds, should be scrutinized to probe the possible involvement of this alternative radical mechanism. In addition to H2, p-block element hydrides also function as reducing agents. While NaBH4/EtOH in combination with CoCl2 or CoBr2 probably promotes a 2-electron cycle,302 the combination of a silane and an H-atom donor (e.g., an alcohol solvent303,304 or 1,4-cyclohexadiene305) seems to promote a radical pathway with Mn or Co catalysts. After initial HAT from L/Mtx+1-H to generate the radical cage 4.22 (Scheme 30), the second H atom is provided by the H-atom donor, as demonstrated by D-labeling experiments. A possible mechanism for the last step, at least when using silane/alcohol as reductant, is proton-coupled electron transfer (PCET) from an intermediate such as 4.23. This final step may also occur in other processes using silane/alcohol as reductant that involve transformations of the HAT-generated radical, such as olefin cross-coupling284–286 or additions to other unsaturated acceptors, to yield the caged product radical 4.24. It should be considered as a viable alternative to the more commonly invoked OSET, which transforms the radical into a carbanion, followed by a final protonation. Curiously, a few of these reduction processes could be improved by the presence of an oxidant (e.g., TBHP). This additive may help by reoxidizing L/Mtx, which accumulates as a result of the unwanted self-terminations. The 1-electron L/Mtx+1-Y (e.g., Y ¼ OtBu) oxidation product can then be reinjected by the silane into the cycle, whereas the silane alone is not able to convert L/Mtx to the active L/Mtx+1-H.

Scheme 30 Possible pathway for radical hydrogenation by PCET.

Reversible Homolysis of Metal-Carbon Bonds

1.03.4.3.2

61

Dehydrometallation

This reaction occurs when the alkyl ligand has at least one H atom in the b position as in 4.25 (Scheme 31) and results in the generation of L/Mtx+1-H and an alkene, 4.26. Many organic transformations incorporate this as an intermediate step or the final one. The ubiquitous b-H elimination, which involves only 2-electron steps (path a) has frequently been invoked. However, only complexes involving homolytically strong metal-carbon bonds have unambiguously been shown to follow this pathway. Bond homolysis followed by b-H atom transfer (path b), i.e., the reverse of the HAT radical generation of Scheme 23B, dominates for homolytically weak bonds. While path a requires a cis-vacant coordination site, path b does not have any specific requirements except for a low Mt-C BDE. Many claims of b-H eliminations for 3d-metal L/Mtx+1-R intermediates should be reconsidered.

Scheme 31 Two pathways for dehydrometallation.

Several cobalt-catalyzed radical reactions yield unsaturated products from a proposed dehydrometalation of alkylcobalt(III) intermediates, with generation of L/CoIII-H. An example is the catalytic olefin isomerization.306 A few isolated L/CoIII-R intermediates were shown to undergo dehydrometallation upon irradiation.268,269 This clearly suggests path b of Scheme 31, because cobalt(III)-C bonds are susceptible to homolytic photocleavage14,307 and the b-H atom abstraction from radicals by L/CoII complexes, which also plays a fundamental role in catalytic chain transfer for polymerization processes,308 is well-documented.

1.03.4.3.3

Alkyl-hydride reductive elimination

If the R%-trapping complex contains a hydride ligand, the L/Mtx+2(H)(R) transient 4.27 may in principle produce R-H by 2-electron reductive elimination (path a in Scheme 32). This requires accessibility of three adjacent oxidation states for the metal atom. If R% is trapped by a metal alkyl complex with a b-H atom on R0 (L/Mtx+1-R0 (bH)), the same alkyl-hydride intermediate may also be obtained by dehydrometallation from a L/Mx+2(R)(R0 (bH)) intermediate 4.28, path b. The occasional proposition of these 2-electron pathways in metal-promoted radical reactions was again inspired by the extensive literature on complexes with homolytically strong bonds. An alternative process, especially if both Mtx+2-R and Mtx+1-H bonds are homolytically weak, is HAT from the L/Mtx+1-H complex to R% (path c). Note that formation of RH via pathway c does not require the presence of a vacant coordination site on L/Mtx+1-H. The anthracene hydrogenation mechanism highlighted in Scheme 20A is an example.254 When a L/Mtx+1-R0 (bH) complex is involved, dehydrometallation (path d) with formation of L/Mtx+1-H, which may be a concerted 2-electron process or an homolysis/HAT sequence (Scheme 31), may occur first followed by path a or c.

Scheme 32 Generation of a saturated product from an alkyl-hydride intermediate. The R0 (bH) symbol indicates a b-H containing alkyl group.

62

Reversible Homolysis of Metal-Carbon Bonds

The mechanistic scenario for the R-H–producing reactions is even more complicated, because the H atom may also be provided by the solvent or other additives with homolytically weak Y–H bonds, especially for reactive radicals. In addition, when R has a b-H atom, R-H may also be obtained by spontaneous disproportionation, together with an equivalent amount of R(-H). The sequence of b-H atom transfer to L/Mtx (Scheme 31) followed by reaction of the resulting hydride complex L/Mtx+1-H with a second radical (Scheme 32) corresponds to a metal-catalyzed disproportionation. Demonstration of which pathway is followed, and even whether the metal complex has any role at all, is not trivial. Metal-promoted or –catalyzed reactions with generation of R-H or mixtures of R-H and R(-H) are abundant, but the studies are seldom accompanied by mechanistic investigations with suitable control experiments, or computations.

1.03.4.3.4

Dialkyl reductive elimination

This step is also frequently invoked in metal-mediated or -catalyzed radical reactions, notably in a growing number of powerful radical cross coupling processes. Like for the alkyl-hydride reductive elimination treated in the previous section, the occurrence of this step requires the accessibility of three adjacent oxidation states (Scheme 33, path a) to generate the dialkyl intermediate 4.29. However, like for the formation of R-H from a hydride intermediate, the formation of R-R0 from an alkyl intermediate may also

Scheme 33 Generation of a cross-coupling product by a radical process.

occur by direct group transfer (Scheme 33, path b), which does not require, like path c of Scheme 32, the presence of a vacant coordination site. The occurrence of the outer-sphere path b, without reversible Mt–R bond formation, has been proposed for Fe-catalyzed Kumada radical cross-couplings involving R-Y and ArMgBr, with [FeIICl2{1,2-[(3,5-R2C6H3)2P]2C6H4}] (R ¼ tBu, SiMe3),309,310 [FeIICl2(IPr)2]311 and [FeIIPh2(IPr2Me2)2]312 (IPr2Me2 ¼ 1,3-diisopropyl-4,5-dimethylimidazol-2-ylidene; IPr ¼ 1,3-diisopropylimidazol-2-ylidene) pre-catalysts. In these reactions, the proposed cycle involves L(Y)/FeII, L(Ar)FeII and L(Ar)FeIII-Y species. Kinetic studies have proven radical escape from the {L(Ar)FeIII-Y,R%} cage, whereas radical rebound within the caged radical pair would yield a putative FeIV species.311 The outer-sphere path b may also be involved in the CuI-catalyzed radical termination (CRT, further commented in Section 1.03.5.11).313 The inner-sphere path a, on the other hand, has been proposed for several other Fe-, Co-, Niand Cu-catalyzed cross-couplings (Kumada, Negishi, Heck, Sonogashira, etc.), though rarely with sufficient supporting evidence. The cycles with most consensus (though not the only ones) involve FeI/FeII/FeIII,314 Co0/CoI/CoII,315 NiI/NiII/NiIII,316 or CuI/CuII/ CuIII species.317 It seems clearly established that Mtx(R0 ) active species prefer to activate alkyl electrophiles (R-Y) by a radical pathway (AT or ET) to generate R% and Y-Mtx+1(R0 ) or [Mtx+1(R0 )]+, followed by radical rebound to yield Y-Mtx+2(R)(R0 ) or [Mtx+2(R) (R0 )]+, whereas aryl electrophiles (Ar-Y) are activated by 2-electron oxidative addition. Either way, the final proposed step is R-R0 or R-Ar or Ar-Ar0 reductive elimination, with R/R0 /Ar/Ar0 coming from either the electrophilic or the nucleophilic reactant. Hence, these FeI/FeIII-, Co0/CoII- or NiI/NiIII-based systems would operate entirely by two-electron elementary steps for ArMgBr reagents, and with (AT or ET)/rebound activation for RMgBr reagents. DFT investigations indicate low barriers for the reductive elimination from FeIII(Ar)(R) and FeIII(Ar)(Ar0 ) species,318,319 but the FeIII-R and FeIII-Ar homolytic bond strengths to probe the alternative one-electron pathway were apparently not assessed.

1.03.4.3.5

Alkyl transfer to an electrophile

For certain metal-catalyzed radical reactions, the diversion toward radical addition to a 2-electron electrophile trap has been taken as evidence of an organometallic intermediate 4.30 (Scheme 34A), in which the nucleophilic character of the C atom is increased. Under these conditions, the metal complex is irreversibly consumed and the process becomes stoichiometric. Suitable control experiments are needed, however, because organometallic intermediates may also be generated by non-radical routes. The N,N-diallyl 2-iodoaniline (4.31) cyclization promoted by MnCl2/nBuMgBr, Scheme 34B,320 is an example. The reaction gave 1-allyl-3-methylindole 4.32 under either stoichiometric or catalytic conditions by probable radical quenching by HAT from the THF solvent, but the addition of E-Y under stoichiometric conditions diverted the reaction toward the 3-ECH2-substituted product 4.33.

Reversible Homolysis of Metal-Carbon Bonds

63

Scheme 34 Trapping of the organometallic intermediate by an electrophile: (A) general scheme; (B) a representative example.

1.03.4.3.6

Oxidation

Certain catalyzed radical reactions on alkene substrates combine reductants and dioxygen. Using NaBH4 under wet conditions, alcohols are obtained as major products, whereas silylperoxides are produced when using a silane under anhydrous conditions. According to the accepted mechanism of Scheme 35A, following HAT activation, the radical reacts with O2 and the generated ROO% collapses to yield a peroxo complex 4.34. This is subsequently hydrolyzed/reduced under wet conditions, or hydrosilylated. The proposed mechanism is supported by studies of individual steps under stoichiometric conditions for a cobaloxime system321,322 and isolated alkyl- and alkylperoxocobalt(III) compounds are active catalysts for these transformations.323 A comparative study of L/Mn, L/Fe and L/Co systems (L ¼ tetra-t-butylphthalocyanine) showed only moderate inhibition by TEMPO for the Co system and a stronger one for those of Mn and Fe.324 Therefore, whether L/Mtx+1-R bond formation to yield 4.35 occurs for Mtx+1 6¼ CoIII is not certain.286 The O2 insertion into the metal-carbon bond via bond homolysis has several precedents for CoIII-R bonds,325–327 whereas direct O2 insertion by an even-electron pathway cannot be easily envisaged for coordinatively saturated compounds.

Scheme 35 Radical-dioxygen interactions: (A) proposed radical mechanism for the oxidation/reduction olefin functionalization; (B) alternative single-oxygen atom insertion.

64

Reversible Homolysis of Metal-Carbon Bonds

The L/CoIII-R system with L ¼ salen and R ¼ CH(O2CMe)CH2Y (Y ¼ poly(vinyl acetate) chain), however, was shown to insert only one O atom by reaction with O2, producing an alkoxy derivative 4.36 in the absence of reducing agents, Scheme 35B.328 This was rationalized by the reversibility of certain bond formations (supported by DFT-calculations of BDEs) and by a kinetic control. After Mtx+1-C bond homolysis for 4.35, both O2 addition and metal alkylperoxide bond formation are reversible, thus residual carbon-based radical can add to the alkylperoxy radical to produce peroxide 4.37, which undergoes reversible O–O bond homolysis and final trapping of the resulting alkoxy radical. This mechanistic variant may also be relevant for certain metal-catalyzed oxidative radical reactions. HAT-generated radicals have also been involved in olefin hydrohydrazination (oxidization by a diazo reagent R0 N ¼ NR0 to yield RR0 N-NHR0 ).329,330 Oxidation by a N-fluoropyridinium salt to a carbocation was proposed to give rise to hydroalkoxylation in an alcohol solvent,331 to intramolecular hydroamination with an internal amine function,332 and to hydrofluorination if no Nu-H substrate is present.333 Intermolecular hydrofunctionalizations with other Nu-H has also been achieved through oxidation by hypervalent iodine compounds (e.g., PhIO)334 or by direct involvement of ArSO2-Nu reagents.330,335–338

1.03.4.3.7

Transmetalation

Transfer of an alkyl radical from one metal to another is a well-established process in preparative organometallic chemistry and in biochemistry151 and may therefore occur if L/Mtx+1-R is formed by radical trapping, including in catalytic cycles. One example is provided by the aldehyde addition to olefins by the cooperative action of L/CoII (L ¼ Salen, catalytic) and CrCl3 (stoichiometric).339 The reaction is proposed to involve radical transfer from an alkylcobalt(III) intermediate 4.38 (Scheme 36), produced by olefin HAT activation, after reduction by CrII to the alkylcobalt(II) intermediate 4.39, to a second CrII ion to yield the alkylchromium(III) complex 4.40 (as documented in stoichiometric processes),86 followed by typical Nozaki-Hiyama-Kishi chemistry with the aldehyde R0 CHO to ultimately produce (R)(R0 )CHOH after work-up. Alkylcobalt(II) derivatives 4.39 appear more susceptible to homolytic bond cleavage than their CoIII counterparts 4.38.340 Hence, this transmetalation from a less to a more electropositive metal imparts carbanion character to the olefin-produced radical. Certain mechanistic details, such as the regeneration of the CoIII-H species, remain obscure. However, the implication of an alkylcobalt(III) intermediate and its reaction with CrII were supported by delayed addition and stoichiometric experiments.

Scheme 36 Transformation of olefins into carbanion equivalents by a radical/polar crossover.

1.03.5

Reversible metal-carbon bond homolysis in controlled radical polymerization

This section briefly describes the principles of the controlled radical polymerization based on reversible metal-carbon bond homolysis in organometallic dormant chains. The contribution of each metal, organized by element type, will then be briefly overviewed.

1.03.5.1

General aspects of organometallic-mediated radical polymerization (OMRP)

Controlled radical polymerization (CRP), also known as Reversible Deactivation Radical Polymerization (RDRP),341 is a preferred tool to produce functional polymers because it gives access to size and dispersity control and to high chain-end functionality, it permits to polymerize a large variety of monomers and is compatible with a wide variety of reaction conditions. Many different CRP methods are available depending on the nature of the moderating agent (T), which must be able to reversibly trap the growing radical chain (Pn%) to yield a Pn-T dormant species. The available methods are grouped into two families, depending on whether the dormant chain is activated by a spontaneous homolytic cleavage (dissociative activation, usually termed “reversible deactivation” polymerization) or is displaced by a degenerative exchange with another radical chain (associative activation, also known as “degenerative transfer” polymerization). In the latter case, the dormant species also plays the role of reversible chain transfer agent. Certain moderating agents are specific for the dissociative activation mode, such as nitroxides (R2N-O%) for Nitroxide-Mediated Polymerization (NMP),342 whereas others are specific for the associative mode, e.g., iodine for Iodine Transfer Polymerization (ITP).343 In this article, we focus on systems where T is a transition metal complex (L/Mtx), yielding an organometallic dormant species (L/Mtx+1-Pn). For this reason, this method has been termed Organometallic-Mediated Radical Polymerization (OMRP).344–346 Depending on the structural characteristics (Mtx+1-C BDE, coordination number and geometry, metal oxidation state and spin state), the metal complex can operate in the dissociative and/or associative mode, Scheme 37.

Reversible Homolysis of Metal-Carbon Bonds

65

Scheme 37 Dissociative and associative activation of the organometallic dormant species in OMRP (M ¼ monomer; ka/kd/kexch/kp/kt ¼ activation/deactivation/ exchange/propagation/termination rate constant; Pn% and Pm% ¼ polymer chains).

In the dissociative activation mode, the number of L/Mtx molecules is at least as high as the number of growing chains and a dynamic equilibrium is established between the dormant chains and the active species plus the moderating agent as discussed in Section 1.03.2.1 (Fig. 1). Consequently, the irreversible self-terminations by coupling and/or disproportionation are minimized and the polymerization control is based on the “persistent radical effect.” The polymer quality (targeted molecular weights, lower dispersity, and higher chain-end fidelity) is improved for homolytically stronger Mtx+1-C bonds, at the expense of slower polymerizations. In the associative activation mode, the free radical concentration exceeds that of the metal complex. In this case, controlled polymer growth requires an open coordination site on the metal center for the entering radical, a much greater rate of degenerative exchange than that of propagation (kexch[L/Mtx+1-R]  kp[M]) and a low but continuous flux of new radicals injected into the solution. The organometallic species acts in this case as a transfer agent, there is no persistent radical effect and the polymerization rate obeys the same law as in free radical polymerization. A low Mtx+1-C BDE is in principle not required in this case, so long as the exchange activation barrier is sufficiently small. An important aspect of transition metal complexes is their possible action as atom transfer catalysts, thus providing a possible interplay of two CRP mechanism. A halogen-terminated polymer chain Pn-Y, where Y is most frequently Cl or Br, may be activated by a L/Mtx complex, yielding L/Mtx+1-Y and the free radical chain, Pn% (Scheme 38, equation a). However, a second L/Mtx molecule may trap the produced radical and generate the organometallic product L/Mtx+1-Pn (equation b). A great many CRP processes have been developed on the basis of the atom transfer activation a, known as ATRP.347,348 ATRP processes do not involve organometallic species so long as the L/Mtx+1-Pn bond is too weak to form. However, the possible intervention of direct radical trapping by the

Scheme 38 One-electron oxidative addition and interplay between ATRP and OMRP controlling mechanisms.

66

Reversible Homolysis of Metal-Carbon Bonds

metal complex in an ATRP process has to be considered.344 While the two trapping mechanisms generally cooperate in maintaining a low steady-state radical concentration and thus improve control, there are also situations where the formation of the organometallic species introduces an unwanted catalyzed radical termination.313 The sum of equations a and b corresponds to a 1-electron oxidative addition, as already highlighted in Section 1.03.4.2.1.1 (Scheme 21).258 Reviews of this area are available,344-346,349–357 some of them dedicated to a specific metal, notably iron354 and cobalt.349,353 An interesting feature of OMRP is the possible BDE modulation by the metal coordination sphere (ligand engineering), specifically in the dissociative activation mode, to adapt the dissociative equilibrium to the reactivity of the radical of interest. The BDE depends on the metal and ligands, but also on the monomer (see below). OMRP has allowed the controlled polymerization of a few challenging monomers, for which other controlling methods have failed.355–357 However, the main drawbacks of OMRP are the need of one metal complex molecule per chain and the need to remove this function by a post treatment for certain applications because of the associated color and toxicity. In ATRP, on the other hand, the L/Mtx complex is catalytic and the recovered polymer has a halogen atom as chain-end functionality. In terms of practical applications of OMRP, particular attention is devoted to L/Mtx complexes that are easily and inexpensively prepared, possibly commercially available, as well as non-toxic (because traces of metal may remain in the polymer even after efficient post-treatments). The monomer scope covers a wide range. The reactivity of a monomer and its associated radical are inversely proportional (Fig. 6) and the propagation rate is dominated by the monomer reactivity. Thus, monomers associated to more stabilized radicals, namely conjugated dienes, vinyl aromatics, (meth)acrylates or acrylonitrile, polymerize faster. Conversely, monomers associated to less stabilized radicals are less reactive. The CRP community often artificially separates the monomers in two families: the “More Activated Monomers” (MAMs) and the “Less Activated Monomers” (LAMs), although the reactivity scale is a continuum. MAMs are easily and successfully polymerized by most controlling methods, whereas most techniques fail for LAMs because they lead to Pn-T with greater BDEs, harder to reactivate. Consequently, most moderating agents that work well for MAMs inhibit polymerization of LAMs. In this respect, the weaknesses of certain L/Mtx+1-C bonds and their tunability by appropriate coordination sphere engineering has placed dissociative OMRP as the top performing method for the CRP of certain LAMs. The associative activation methods are in principle not affected by the Pn-T bond strength and thus are potentially more promising for LAMs. However, an additional important limitation is related to the monomer addition regioselectivity. For asymmetric monomers, the growing radical chain can add to both the less substituted C atom (tail) to produce a head radical, Pn,H%, e.g., PVC-CH2-CHCl% for VC, or the more substituted C atom (head) to produce a tail radical Pn,T%, e.g., PVC-CHCl-CH2%. Whereas asymmetric MAMs display a 100% regioselectivity, because Pn,H% is much more stabilized and is produced via a lower activation barrier, head and tail radicals have closer stability and addition barriers for LAMs. This leads to a significant fraction of inverted monomer additions during the propagation step (e.g., around 1–2% for PVAc or 4–5% for PVDF). Chain propagation appears not to be affected by cage effects, thus all parameters related to propagation (regioselectivity, tacticity, copolymerization statistics, etc.) do not depend on the trapping agent. The tail dormant species generated after an inverted monomer addition has a stronger Pn,T-T bond, hence reactivation is slower for both dissociative and associative activation methods and this species progressively accumulates. For dissociative methods, this leads to a polymerization slowdown (or full stop). For the associative ones, the rate depends on the continuous injection of new radicals into the solution, hence polymerization does not stop but there is a slowdown (or full stop) of the radical exchange process, which is responsible for control. Therefore, increasing dispersities and loss of chain-end functionality are observed, tending toward the situation of a free radical process. For certain monomers, OMRP has been able to provide ad hoc solutions by equalizing the reactivation barriers of head and tail dormant species through monomer chelation or bond polarity effects.357

1.03.5.2

Titanium

Compounds with a TiIV-R bond are extensively used to polymerize alkenes by the coordination/insertion mechanism. Nevertheless, this bond has a sufficiently low homolytic strength for reversible homolysis in certain systems and thus they can mediate radical polymerizations. Since 2004,359 Asandei and coworkers have applied this chemistry for the controlled polymerization of

Fig. 6 Reactivity scale of MAMs and LAMs.358 BD ¼ 1,3-butadiene; AN ¼ acrylonitrile; MMA ¼ methyl methacrylate; St ¼ styrene; VC ¼ vinyl chloride; VAc ¼ vinyl acetate; VDF ¼ vinylidene fluoride; E ¼ ethylene.

Reversible Homolysis of Metal-Carbon Bonds

67

Fig. 7 Titanium(IV) complexes tested in radical polymerization.

styrene,359–373 butadiene374–378 and isoprene.379,380 The mediating agent is a TiIII complex, generated in situ by reduction of a stable TiIV precursor. Zinc was shown to be the best reducing agent364 and [Cp2TiCl2] (5.1, Fig. 7, Y ¼ Cl) the best precursor among several investigated TiIV complexes,362-364,381,382 yielding [Cp2TiIIICl]. This complex operates, at the same time, as polymerization moderator and as part of the initiating system in combination with various other molecules, see Scheme 39. The [Cp2TiIIICl] reducing power, combined with the oxophilicity/halidophilicity of titanium(IV), produces primary radicals by electron transfer in combination with many stable molecules such as aldehydes,361,383 active halides R-Y (a,a0 -dihalo-p-xylene371,375,377,378, 1,10-dibromodecane,373 4-methoxybenzenesulfonyl chloride,368 (1-bromoethyl)benzene369), peroxides360,365,366,384 and also epoxides by radical ring opening (RRO).359,364,367,372,376,380 Subsequently, a second [Cp2TiIIICl] molecule reversibly traps the growing radical chain to generate the organometallic TiIV dormant species, rather than Cl atom transfer to generate a Cl-capped chain (as in reverse ATRP, i.e. equation a in Scheme 38). The NMR and IR analyzes of the recovered polymer products obtained by aldehyde or epoxide RRO initiations have confirmed the presence of the Cp2TiCl-O fragment at the a chain-end and the absence of Cl atoms at the o

Scheme 39 Initiation systems and moderating equilibrium for the [Cp2TiIIICl]-mediated OMRP.

chain-end (hence, ruling out an ARTP mechanism).359,375,377,380 Fig. 7 summarizes the Ti complexes that have been tested as moderator or initiator/moderator. Among the sandwich complexes (5.1–5.4) with different Z5 and monodentate wedge ligands,362-364,381,382 the most efficient one in terms of controllability is the

68

Reversible Homolysis of Metal-Carbon Bonds

simplest and least expensive [Cp2TiCl2].359 The half-sandwich complexes (5.5–5.7) required higher temperatures for TiIV-C bond homolysis and also led to broader dispersities.363 Improved control resulted from a decrease of the five-membered ring electron donating power (tBuCp, IPrCp, EtCp and Cp).363 The alkoxide (5.8) and diketonato (5.9–5.10) complexes were tested in styrene polymerization.381,385 Only 5.8 (Y ¼ Cl, OiPr) provided some control, although with low initiation efficiencies, whereas 5.9 only led to free radical polymerization and the TiIII complex 5.10 gave no initiation. The readily accessible aminebis(phenolate) complexes of TiIII, 5.11, provided excellent control (Đ ¼ 1.14) for the MMA polymerization in toluene at 80  C in combination with azo-initiators (V-70, AIBN).386 This polymerization, however, does not involve reversible TiIV-C homolysis. Rather, after radical initiation, the growing polymer chains are trapped as enolates (TiIV–O bond) and continue to grow by a bimetallic Group Transfer Polymerization (GTP) mechanism. In this respect, other reports of MMA polymerizations with radical initiators in the presence of highly oxophilic metal complexes, including [Cp2TiCl2],387–389 and described as controlled radical polymerizations, may need reconsideration. The OMRP mechanism with this highly oxophilic metal is unambiguously demonstrated only for styrene and diene monomers. An attempt to polymerize VDF with this system led to no polymer production.390 The effect of various parameters (solvent and additives,382 metal/initiator/monomer stoichiometry, temperature,391 and initiation method370,372) has been studied in detail for the best-performing [Cp2TiIIICl] system. Initiation by aldehyde reduction, yielding Cp2ClTiIV-OCHR%, is more efficient than the RRO method (lower proportion of side reactions). Peroxides may function either by the classical thermal decomposition or by a redox process and the resulting ether chain-end groups are less easily postmodified. The alkyl halide initiation method is the most sensitive and difficult one to optimize.372 Low-dispersity (1.10–1.18) polystyrenes were obtained with each initiation modes, the optimum temperature being 70–90  C for the [Cp2TiIVCl2]/Zn system.370-372,392 For polyisoprene, a lower but still satisfactory level of control (Đ of 1.2–1.3) and optimum initiator efficiency were achieved for [DBPX]/[Cp2TiCl2]/[Zn] ¼ 1/6/20 at 70  C.379 Random and block styrene-isoprene copolymers with Đ of 1.39–1.50 were also obtained using [Cp2TiCl2] at 110  C.393 Polybutadiene, on the other hand, could only be obtained with Đ  1.5.376

1.03.5.3

Vanadium

Bis(iminopyridine)vanadium(III) complexes 5.12, also active coordination/insertion polymerization catalysts for olefins and dienes,394–397 are able to control the radical VAc polymerization (Đ ca. 1.3) at 120  C using AIBN as thermal initiator (Scheme 40). Control, however, was satisfactory only for relatively small monomer conversions (ca. 30%) because of slow

Scheme 40 Bis(iminopyridine) vanadium complexes used in OMRP and their mechanism of action.

moderating complex decomposition, possibly by irreversible radical additions to the non-innocent diiminopyridine ligand.398,399 The main point of interest is the controlling mechanism, because the presence of Cl ligands raises the question of a possible reverse ATRP via the reversibility of the reduction step to 5.13. Additionally, OMRP may be promoted by either a VIII/VIV-R pair (5.12 with a putative VIV radical trapping product) or by the VII/VIII 5.13/5.14 pair generated after in situ Cl atom transfer. The

Reversible Homolysis of Metal-Carbon Bonds

69

presence of metallic chain-ends was indicated by the isolated polymer color, 1H NMR and derivatization studies. A distinction between VIII and VIV at the chain end was possible from the polymer molecular weights, because direct trapping by VIII would produce one chain per initiating radical, whereas reduction to VII consumes two radicals per vanadium—one to initiate the propagation and one for metal reduction—hence double average molar masses. Further characterization by EPR and XPS of the recovered polymer, the identification of Me2C(CN)Cl in solution, and validation by DFT calculations confirmed the mechanism shown in Scheme 40.398,399 Extension of this system to other vinyl esters (propionate, pivalate) gave equally good control, whereas vinyl benzoate and styrene are less well-controlled (Đ > 1.5)398,399 and other more activated monomers (methyl methacrylate, acrylonitrile) showed no control.399 Therefore, the VIII-C BDE for this system is only suitable for more reactive radicals. Reducing the steric bulk of the aryl ortho-substituents or removing them altogether (i.e., R2 ¼ C6H3-2,6-Et2, C6H3-2,6-Me2, Ph) led to poorer control, whereas electronic variations at the para position are relatively unimportant.400 Steric bulk may offer protection against the unwanted ligand attack by the radical, without affecting the productive reversible addition to the metal center. Use of a bulky aliphatic R2 group (Cy) also led to poor control, with fast uncontrolled polymerization suggesting a weakening of the VIII-C bond. Varying the R1 substituents while keeping the same R2 (C6H3-2,6-iPr2) confirmed the steric protection hypothesis: for H, rapid complex degradation only led to short oligomers, but good control was maintained for R1 ¼ Et and iPr. Other vanadium complexes (Fig. 8) were less efficient.399 Polystyrene radical chains are apparently incapable of reducing VIII-Cl to VII and/or the added (or in situ produced) VII complex does not efficiently trap the chains. The more strongly binding PVAc% chains, on the other hand, produce polymers with a certain degree of control, especially when used in combination with the VIII

Fig. 8 Other vanadium complexes tested in OMRP.

precursors (5.19 and 5.20), but lower than with the above-described bis(imino)pyridine systems 5.12.

1.03.5.4

Chromium

Starting in 1978, Minoura et al. described the polymerization of various vinyl monomers by redox initiation with the combination of chromous acetate and a peroxide and pointed out their “living” character.401–404 Controlled polymerizations were also obtained when using an “aged” system, namely by introducing the monomer after all Cr2+ was converted to Cr3+, although longer aging gave lower initiation efficiencies. The polymerizations were kinetically well-defined, with slower monomer consumption than in the free radical process and a linear molar mass increase with monomer conversion. They were controlled, however, only at low temperatures ( Xyl > Dipp; methods A and B).

1.03.5.5

Molybdenum

The 17-electron MoIII complexes 5.26–5.28 (Fig. 10) control the AIBN-initiated styrene polymerization and are ineffective for the MMA polymerization.411 The absence of halogenated chain ends excludes an ATRP mechanism and suggests the formation of organomolybdenum(IV) dormant species. The same complexes, however, are also ATRP catalysts with bromoalkane initiators through reversible Br-atom transfer with production of a L/MoIV-Br moderator. Obviously, the L/MoIV-Pn dormant species may also form under ATRP conditions and the two moderating equilibria cooperate, as shown in Scheme 38.

Fig. 10 Molybdenum-based complexes used in ATRP/OMRP.

The RN]CHdCH]NR (R2-dad) ligands gave access to steric modulation for complexes 5.29, the structures of which are more consistent with a MoV ene-diamido rather than a MoIII(diazadiene) formulation. The sterically more encumbered 5.29a led to reversible dissociative activation of polystyrene chains,412 whereas the less congested 5.29b gave irreversible trapping of both polystyrene and poly(methyl acrylate) chains. Moving from to the diiodo complex 5.29c, however, restored reversible homolysis.413 Octahedral [MoIIIY3(PMe3)3] complexes (5.30) efficiently control styrene polymerization by ATRP,57,414 but poorly perform by OMRP (thermal AIBN initiation). With 5.30a, slower polymerization than without metal complex and a polymer molar mass increase with conversion suggest the presence of a moderating equilibrium, though insufficient for good control. Conversely, identical results as in free radical polymerization were obtained with 5.30c.57 Complex 5.31 was also tested as a unimolecular initiator, probing a possible MoIII-CH2SiMe3 bond homolysis, but no polymerization occurred at 80  C. Polystyrene did form at 110  C, but at similar rate and with the same characteristics as for the self-initiated metal-free process. This suggests a significant Mo-C BDE decrease upon increasing the metal oxidation state from III to IV, as also indicated by DFT calculations. It also suggests that the self-initiated polystyrene radical chains are not efficiently trapped to yield a putative CpMoIV(Pn)(CH2SiMe3)2, underlining the importance of the one-electron ligand (Cl in 5.28 vs. CH2SiMe3 in 5.31).411

72

Reversible Homolysis of Metal-Carbon Bonds

1.03.5.6

Manganese and rhenium

Although homolytically weak Mn–C bonds for typical coordination compounds (oxidation state II and above) seem to play a role in metal-mediated/catalyzed organic reactions (Section 1.03.4), their contribution in controlled polymerization processes is undocumented. The same is true for Re–C bonds. The [MtI(CO)5%] radicals (Mt ¼ Mn, Re), generated from the photolytic splitting of their dimers, are able to initiate the polymerization of a few monomers. The resulting polymers feature [(CO)5Mt-Pn] chain ends, indicating the initiation ability of the metalloradicals, but no moderating action was demonstrated for these polymerizations.415 The same radicals were subsequently used in the CRP of VDF by the iodine transfer (ITP) controlling method (Scheme 43).416,417 However, the role of these radicals appeared limited to the reactivation by iodine atom abstraction of the less reactive iodine-capped dormant species, which forms by non-degenerative transfer after an inverted (head-to-head) monomer addition. The formation and possible reactivation of the organometallic dormant species, [(CO)5Mt-PVDFH/T], which may also be produced by the direct trapping of the growing chain by the continuously generated [Mt(CO)5%] radicals, was not considered.

Scheme 43 Action of the [Mt(CO)5%] (Mt ¼ Mn, Re) radicals in the reactivation of PVDFT-I dormant species in the ITP of VDF and possible role of metal-carbon bond homolysis.

For this reason, subsequent work has analyzed the homolysis of model complexes of these putative dormant species for manganese, namely [(CO)5Mn-RF] (RF ¼ CF3, CHF2, CH2CF3, 5.32 in Fig. 11), as well as [(CO)5Mn-COCF2CH3] (5.33),60 finding that these bonds are too strong (46–54 kcal mol−1, Table 1) for significant generation of RF% under the VDF ITP conditions. However, they photoinitiate the VDF polymerization under both visible and UV light irradiation, though without ensuring controlled growth.418 Therefore, if any such [(CO)5Mt-PVDFH/T] dormant species form during the ITP process, they can be reactivated under irradiation but the controlled polymer growth is entirely ensured by ITP. The significantly weaker MnI–C bond (35.3  2.8 kcal mol−1) in [(CO)5Mn-CH(Me)COOMe] (5.34) allows thermal initiation of methyl acrylate polymerization at relatively low temperatures, but the polymerization was uncontrolled, as expected from the rapid [Mn(CO)5%] radical disappearance by dimerization.73

Fig. 11 Half-sandwich manganese complexes used in radical polymerization.

Reversible Homolysis of Metal-Carbon Bonds

73

A few additional half-sandwich MnI (18-electron) and MnII (17-electron) complexes (compounds 5.35–5.38 in Fig. 11) were tested as moderating agents for the polymerization of MMA initiated by AIBN,419 but polymerization rates and polymer properties similar to the metal-free control were observed in all cases, indicating no significant moderating effect. When heated in the presence of AIBN and in the absence of monomer, the 18-electron 5.35 and 5.36 decomposed, whereas 5.37 and 5.38 were stable. The latter could in principle bind an organic radical but the product would be an unlikely organomanganese(III) complex. Indeed, DFT calculations indicated that only a weak van der Waals adduct is formed between the model PhCH%Me radical and 5.37, with an insignificant stabilization energy of −1.4 kJ mol−1 (ca. −0.3 kcal mol−1).

1.03.5.7

Iron

The area of iron-mediated CRP (both ATRP and OMRP) has been reviewed in 2014.354 Although there is intensive activity on Fe-based ATRP catalysts, only few contributions address the possibility of reversible Fe–C bond homolysis. The main investigated systems are shown in Fig. 12. Under thermal AIBN initiation, the porphyrin (5.39), phthalocyanine (5.40) and Schiff base (5.41, 5.42) systems can control styrene polymerization, although with higher than targeted molar masses and rather high dispersities.420 While VAc polymerization was inhibited by 5.40, controlled growth could be provided by complex 5.43, demonstrating an FeIII-C BDE tuning by the coordination environment (weaker for O4 relative to O2N2 and N4).421 For VAc, however, 5.43 performed more poorly than its cobalt analog (Section 1.03.5.9) because of a lower trapping efficiency and the formation of [FeII(acac)2] oligomers, reducing the moderator efficiency. Control improved when operating under more dilute conditions or in the presence of Lewis bases, especially PMe2Ph. Complex 5.43 also controlled the VAc polymerization by degenerative transfer (Scheme 37), indicating its ability to undergo associative radical exchange. The isolated [FeIII(acac)2-PVAc] was shown to function as a single-component macroinitiator for the OMRP of VAc. Similar results were also reported when using [FeIII(acac)3] in the presence of a reducing agent (e.g., ascorbic acid).422 The a-diimine system 5.44 in combination with alkyl halide initiators controls styrene polymerization by ATRP. Although the involvement of direct radical trapping (OMRP/ATRP interplay) was invoked as a path leading to catalytic chain transfer via b-H elimination,423 a later study suggested that this occurs by HAT instead, without FeIII–C bond formation.424 The bis(phenolate) systems 5.45, on the other hand, lead to ATRP/OMRP interplay for the polymerization of substituted styrenic monomers and MMA after in situ reduction (reverse ATRP conditions).425–427 The reduced FeII complex 5.46 (presumably dimeric in the solid state) was isolated and independently shown to exert control in the V-70-initiated polymerization of styrene and MMA. For VAc, on the other hand, irreversible radical trapping occurs.428 DFT calculations suggested that the diaminebis(phenolate) ligand geometry does not allow the FeII system to attain the preferred tetrahedral geometry whereas the trigonal bipyramidal dormant species is not strained, thus leading to a stronger FeIII–C bond relative to system 5.44.95

Fig. 12 Main iron complexes used in OMRP.

74

Reversible Homolysis of Metal-Carbon Bonds

1.03.5.8

Ruthenium and osmium

Ruthenium yields very successful ATRP catalysts of type [RuIICl2L3]429 and [Cp RuIIYL2],430 working through the RuII/RuIII-Y couple, but radical trapping to yield RuIII-capped dormant chains has not been demonstrated. The osmium complex [OsIICl2(PPh3)3], on the other hand, contributes with reversible chain trapping to the alkyl halide-initiated ATRP of styrene and (meth)acrylates. This is shown by its independent ability to provide moderate control in the AIBN-initiated polymerization.431 The vertical trend in Mt-R BDE (Ru < Os) is responsible for this phenomenon, as confirmed by DFT calculations.432

1.03.5.9

Cobalt

Cobalt is the most investigated and successful metal for OMRP.349,353 It always operates via a diamagnetic 5- or 6-coordinate [L/ CoIII-Pn] dormant species and the [L/CoII] moderator may be a 4-coordinate (square planar or tetrahedral) or 5-coordinate complex, either spin doublet or quartet. In the presence of donor solvents or additives (D), 4-coordinate [L/CoII] may also afford 6-cordinate (18-electron) [L(D)/CoIII-Pn]. The alkylcobalt(III) compounds 5.47 (R ¼ CH2tBu)433 and 5.48434,435 (Fig. 13) were first used as initiators in a dissociative activation mode, yielding controlled polymerizations of acrylates under thermal or photochemical activation, respectively. The polymerizations can also be initiated by a classical radical source with a cobalt(II) complex. In this case, injection of a substoichiometric amount of primary radicals (R0%) can only sustain the dissociative activation mode, whereas an excess (R0%/CoII > 1) may ensure control by DT (Scheme 37, provided a vacant site is available in [L/CoIII-R] for an associative exchange.

1.03.5.9.1

Porphyrin systems

After the seminal report of OMRP with 5.47, further studies have focused on more readily accessible CoII species (Fig. 14) and a classical radical source. The lipophilic 5.49a-c control the polymerization of acrylate esters,436–440 while water-soluble 5.49d,e control the polymerization of acrylic acid in water.441 The intermediate polarity of 5.50 is compatible with a wider array of lipo/ hydrophilic acrylates and acrylamides in both polar and non-polar media and also allows to control tBA, whereas 5.49a is inefficient.442 This results from the CoIII-PtBA bond strengthening by the release of steric strain linked to the removal of two

Fig. 13 First cobalt complexes used in OMRP.

Fig. 14 Cobalt porphyrin complexes used in OMRP.

Reversible Homolysis of Metal-Carbon Bonds

75

o,o0 -Me groups from one aryl substituent, as confirmed by DFT calculations.443 The dissociative OMRP of VAc with 5.49a is completely inhibited, but becomes possible by DT, though control is good only up to moderate conversion.444–446 Dissociative activation of [(TMP)CoIII-PVAc] becomes possible by addition of excess pyridine,446 by the same principle that was elucidated first for [CoII(acac)2] (see next section Scheme 44).447

Scheme 44 Effect of donor additives in the [Co(acac)2]-mediated polymerization of vinyl esters and amides.

After thermal AIBN-initiated MA or DMA polymerization with 5.49a and 5.50, the resulting [L/CoIII-PMA] and [L/CoIII-PDMA] (Mn ¼ 1–1.8  104 g mol−1) were used as macroinitiators for the room temperature OMRP of acrylamides under visible-light irradiation, demonstrating a positive effect of photocleavage in dissociative OMRP. Well-controlled diblock copolymers were obtained only with moderate light intensity.448,449 Complex [(TMP)CoIII-CO2Me] also photoinitiates the OMRP of acrylates and acrylamides at room temperature, with good control even for tBA.449

1.03.5.9.2

b-Diketonate systems

The application of b-diketonate and related systems (Fig. 15) has marked a turning point in OMRP with cobalt. The O4 donor set gives a weaker CoIII-C bond than the N4 donor set of porphyrins, making these systems better suited for the OMRP of LAMs. The simplest and commercially available 5.51a affords excellent control for VAc in bulk and aqueous suspension, with either thermal (V-70 at 30–40  C)450,451 or redox (benzoyl or lauroyl peroxide)452 initiation. 5.51b453 and 5.52454 yield similar results, whereas 5.51c is inefficient. Benzene ring fusion in the salicylate (5.53)455 and 9-oxyphenalenone (5.54)456 systems results in poorer control, while 5.54 also leads to CCT, which is rarely observed for LAMs. 5.51a controls well also other LAMs such as other vinyl esters,457 vinyl amides,457–461 vinyl chloride,462 VDF,463,464 and the copolymerizations of VAc with ethylene,465 1-octene466,467 and perfluorohexyl ethylene.468 More reactive monomers are not well controlled by 5.51a, but acceptable results were obtained for acrylonitrile469,470 and nBA471 under optimized conditions. Allyl radicals, being quite stabilized, cannot be efficiently trapped by 5.51a. Thus, addition of dienes to dormant [L/CoIII-Pn] chains results in the formation of Pn-diene% which, because of the slow diene radical homopropagation and termination by coupling, selectively generates mid-chain-functionalized Pn-diene-dienePn products, including symmetric AnB2mAn triblock copolymers from [(acac)2CoIII-BmAn] diblock dormant chains.472–474 Use of the unimolecular [(acac)2CoIII-(VAc)nR0] initiator (n  4, R0 ¼ Me2C(OMe)CH2CMeCN) in the dissociative mode generally leads to better control. This compound is prepared by V-70/5.51a-initiated VAc polymerization at small VAc/5.51a ratio and is sufficiently stable to be purified by chromatography.475

76

Reversible Homolysis of Metal-Carbon Bonds

Fig. 15 b-Diketonate and cobalt complexes used in OMRP.

Donor solvents or additives (D) strongly affect the polymerization rate and control mode with 5.51a.447 In their absence, the 5-coordinate dormant species can be stabilized by carbonyl group chelation from the chain ultimate monomer,475 if this is a vinyl ester or amide (Scheme 44). The chelation equilibrium still allows access to the vacant site for a DT polymerization. Coordination of D, on the other hand, blocks DT polymerization and stabilizes the moderating species, increasing the propensity to dissociative activation. The latter effect is modulated by the D concentration and binding constant, water being a particularly efficient donor.476 Since water is a common contaminant of commercial 5.51a and solvents, it is possible to witness dissociative activation or shorter than expected induction times for OMRP-DT under supposedly “anhydrous” conditions.477 Chelation by the ultimate monomer rationalizes a few observations such as the reactivity trend for N-vinyl lactams of different ring size,460 the lack of control for g-methylene-g-butyrolactone (though its copolymerization with VAc is controlled)478 and the absence of slowdown or of a worsening of control at high conversions for the VAc polymerization, contrary to other [L/Mtx+1-PVAc] species where this chelation is impossible.143 The latter phenomenon results from compensation of the stronger CoIII-CH2CH(OAc) bond by a weaker 6-membered chelate in the tail dormant species formed after an inverted monomer addition, relative to the weaker CoIII-CH(OAc)CH2 bond and stronger 5-member chelate in the head dormant species. VDF, like VAc, is better controlled by OMRP with 5.51a than by other techniques. This is not due to a chelation phenomenon but rather to the polar effect of the a- and b-F substituents on the BDE, as shown by DFT calculations. Thus, the [(acac)2CoIII-CF2CH2-PVDF] (head) and [(acac)2CoIIICH2CF2-PVDF] (tail) dormant species have bonds of equal strengths, whereas the other techniques, after an inverted monomer addition, yield less labile tail dormant species.120 Steric effects also modulate the [L/CoIII-Pn] BDE: the VAc OMRP-DT is equally fast with 5.51a and 5.51b, but much slower with the former in the dissociative mode without D. Conversely and for the same steric reason, the D effect is stronger for 5.51a.453

1.03.5.9.3

Other planar macrocyclic systems

The success of cobalt porphyrins has naturally led to interest in other related systems with well-developed [L/CoIII-R] chemistry. After the seminal report of the acrylate polymerization with 5.48,434,435 the water-soluble cobaloxime has not led to other OMRP applications, whereas it has extensively been investigated as a CCT catalyst, mostly for methacrylates and styrenics.479–482 The equally water-soluble cobalamin has yielded the controlled polymerization of 2-hydroxyethyl acrylate at pH 7, whereas polyethylene glycol methacrylate gave catalytic chain transfer oligomerization.483 Greater attention has been devoted to salicylidene diamine (Salen-type) systems (Fig. 16). A moderately controlled MA polymerization occurs when initiated by V-70/5.55 at 50  C,484 slowly for R0%/CoII < 1 by dissociative activation and rapidly by OMRP-DT with excess radicals. When using [(Salen)CoIII-Et] as initiator, the strong CoIII–Et bond relative to CoIII-PMA results in poor initiation efficiency. The substituted trans-cyclohexane-1,2-diyl-bridged (Salen ) system 5.56a controls both MA and VAc with AIBN initiation at 60  C. While both dissociative and associative activation modes cooperate for MA, VAc polymerizes only by OMRP-DT. The dissociation equilibrium constants for [(SalentBu,tBu)CoIII-PMA] and [(SalentBu,tBu)CoIII-PVAc] were estimated as 4.2  10−8 and tBu (14.1) > OMe (13.7) > NMe2 (13.3) (values are BDEs in kcal mol−1). For VAc at 120  C, 5.56b gave low conversions without any control, while higher conversions and Mn close to the theoretical values were given by the other systems. On the basis of the above-mentioned studies,485,487 the intervention of OMRP-DT with excess radicals seems possible, at least for 5.56a. However, the generation of better-controlled PVAc when the AIBN/5.56 ratio was lowered to 0.4 for 5.56c suggests that genuine dissociative activation may also occurs. The styrene polymerizations were only poorly controlled and those of MMA only gave CCT oligomers.490

1.03.5.9.4

Other ligand systems

The additional coordination geometries shown in Fig. 17 have also been considered for OMRP moderating agents. The 1,3-bis (2-pyridylimino)isoindolates (bpi) complexes 5.57 provide control for MA and nBA under thermal initiation (V-70/CoII ¼ 1:1) at 60  C. DFT calculations show a negligible effect of the bpi substituents on the [(bpi)(acac)CoIII-CH(CO2Me)Et] BDE, in agreement with the polymerization results.491 Systems 5.58 control styrene and MMA only rather poorly under ATRP conditions (ethyl-2-bromo-isobutyrate initiator). OMRP trapping may be involved in these polymerizations, but the polymer tacticity suggests that non-radical mechanisms may also contribute for 5.58a,b.492 The b-ketoiminates 5.59 are isolobal with the b-ketonates by substitution of one O atom with NR and thus may sterically tune the CoIII–Pn bond strength. Similarly to 5.51a, these systems are able to promote the VAc dissociative and associative OMRP, but slower radical trapping (in the order c > a > b) leads to poorer control and D additives have little impact on the polymerization results.493 The bis(phenoxy-imine) complexes 5.60 also show poor trapping ability in the thermally initiated VAc polymerization. In addition, 5.60b,d,e,f revealed CCT activity. Excellent control could be achieved for VAc with 5.60a when photoinitiated by ArCOP(O)Ph2 (Ar ¼ 2,4,6-C6H2Me3) at 24  C, whereas MA polymerization could not be controlled under any conditions.494

Fig. 17 Other coordination spheres for cobalt complexes used in OMRP.

78

Reversible Homolysis of Metal-Carbon Bonds

1.03.5.10 Rhodium The (octaethylporphyrinato)rhodium(II) dimer, [(OEP)RhII]2, adds irreversibly to both head and tail ends of acrylates to yield [(OEP)RhIII-CH2CH(CO2R)-RhIII(OEP)]. Conversely, the steric demand of the mesityl substituents precludes both dimerization and strong binding to the acrylate head end for the mononuclear [(TMP)RhII] complex. However, tight binding to the acrylate tail end is allowed, yielding [(TMP)RhIII-CH2CH(CO2R)CH(CO2R)CH2-RhIII(TMP)] after dimerization by head-head coupling. The RhIII–CH2 bond is too strong for reversible dissociation, but the photoinitiated MA polymerization is controlled, demonstrating moderation through reversible binding to the head end.495 The high cost of rhodium discourages further development of OMRP applications.

1.03.5.11 Copper Copper is undoubtedly the most successful metal for ATRP through [L/CuI] and [L/CuII-Y] complexes as catalyst and moderator, respectively, but there are no reports of successful [L/CuI]-based OMRP. The ATRP activity is very ligand dependent, spanning several orders or magnitude.124,496 Interplay with OMRP for these systems (Scheme 38) would lead to [L/CuII-Pn] dormant chains. Organocopper(II) compounds are rare and characterized by low BDEs,497 hence suggesting that [L/CuI] complexes might be suitable as OMRP moderators for LAMs. Indeed, reversible radical trapping by [L/CuI] has been evidenced by a slowdown of the polymerization rates.498–500 However, the resulting [L/CuII-Pn] bond is apparently too weak to give sufficient organometallic species to control chain growth. Another problem negatively affecting OMRP with [L/CuI] complexes is catalyzed radical termination (CRT), which has been highlighted for acrylate polymerizations.313,499–504 This occurs via a [L/CuII-Pn] intermediate, which then reacts with a second radical to yield terminated chains and regenerate [L/CuI]. The most active ATRP catalysts, which also lead to the strongest [L/CuII-Pn] bonds,130,500 are also the most efficient radical termination catalysts. The negative impact of CRT in an ATRP process can be minimized by working at very low catalyst concentrations, because CRT depends on [L/CuI], whereas the ATRP activity depends only on the [L/CuI]/[L/CuII-Y] ratio.505,506

Acknowledgment We are grateful to the CNRS for continued support of our research and to the ANR (grant POLYSWITCH No. ANR-19-CE07-003101) for the PhD fellowship of M.M.

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. 29. 30. 31. 32.

Seyferth, D. Organometallics 2001, 20, 2940–2955. Grignard, V. C. R. Hebd. Seances Acad. Sci. 1900, 130, 1322–1324. Synfacts 2018, 14, A155–A159. Gilman, H.; Jones, R. G.; Woods, L. A. J. Org. Chem. 1952, 17, 1630–1634. Halpern, J. Polyhedron 1988, 7, 1483–1490. Halpern, J.; Maher, J. P. J. Am. Chem. Soc. 1965, 87, 5361–5366. Chock, P. B.; Halpern, J. J. Am. Chem. Soc. 1969, 91, 582–588. Endicott, J. F.; Ferraudi, G. J. J. Am. Chem. Soc. 1977, 99, 243–245. Halpern, J. Acc. Chem. Res. 1982, 15, 238–244. Pratt, J. M. J. Chem. Soc. 1964, 5154–5160. Schrauzer, G. N.; Sibert, J. W.; Windgassen, R. J. J. Am. Chem. Soc. 1968, 90, 6681–6688. Pratt, J. M.; Whitear, B. R. D. J. Chem. Soc. A 1971, 252–255. Zou, X.; Brown, K. L.; Vaughn, C. Inorg. Chem. 1992, 31, 1552–1554. Brown, K. L.; Zou, X. Inorg. Chem. 1992, 31, 2541–2547. Chen, E. F.; Chance, M. R. J. Biol. Chem. 1990, 265, 12987–12994. Chen, E.; Chance, M. R. Biochemistry 1993, 32, 1480–1487. Walker, L. A.; Jarrett, J. T.; Anderson, N. A.; Pullen, S. H.; Matthews, R. G.; Sension, R. J. J. Am. Chem. Soc. 1998, 120, 3597–3603. Walker, L. A.; Shiang, J. J.; Anderson, N. A.; Pullen, S. H.; Sension, R. J. J. Am. Chem. Soc. 1998, 120, 7286–7292. Shiang, J. J.; Walker, L. A.; Anderson, N. A.; Cole, A. G.; Sension, R. J. J. Phys. Chem. B 1999, 103, 10532–10539. Yoder, L. M.; Cole, A. G.; Walker, L. A.; Sension, R. J. J. Phys. Chem. B 2001, 105, 12180–12188. Cole, A. G.; Yoder, L. M.; Shiang, J. J.; Anderson, N. A.; Walker, L. A.; Banaszak Holl, M. M.; Sension, R. J. J. Am. Chem. Soc. 2002, 124, 434–441. Sension, R. J.; Cole, A. G.; Harris, A. D.; Fox, C. C.; Woodbury, N. W.; Lin, S.; Marsh, E. N. G. J. Am. Chem. Soc. 2004, 126, 1598–1599. Sension, R. J.; Harris, D. A.; Cole, A. G. J. Phys. Chem. B 2005, 109, 21954–21962. Sension, R. J.; Harris, D. A.; Stickrath, A.; Cole, A. G.; Fox, C. C.; Marsh, E. N. G. J. Phys. Chem. B 2005, 109, 18146–18152. Shiang, J. J.; Cole, A. G.; Sension, R. J.; Hang, K.; Weng, Y. X.; Trommel, J. S.; Marzilli, L. G.; Lian, T. Q. J. Am. Chem. Soc. 2006, 128, 801–808. Jones, A. R.; Hay, S.; Woodward, J. R.; Scrutton, N. S. J. Am. Chem. Soc. 2007, 129, 15718–15727. Harris, D. A.; Stickrath, A. B.; Carroll, E. C.; Sension, R. J. J. Am. Chem. Soc. 2007, 129, 7578–7585. Robertson, W. D.; Warncke, K. Biochemistry 2009, 48, 140–147. Jones, A. R.; Woodward, J. R.; Scrutton, N. S. J. Am. Chem. Soc. 2009, 131, 17246–17253. Stickrath, A. B.; Carroll, E. C.; Dai, X.; Harris, D. A.; Rury, A.; Smith, B.; Tang, K.-C.; Wert, J.; Sension, R. J. J. Phys. Chem. A 2009, 113, 8513–8522. Peng, J.; Tang, K.-C.; McLoughlin, K.; Yang, Y.; Forgach, D.; Sension, R. J. J. Phys. Chem. B 2010, 114, 12398–12405. Jones, A. R.; Hardman, S. J. O.; Hay, S.; Scrutton, N. S. Angew. Chem. Int. Ed. 2011, 50, 10843–10846.

Reversible Homolysis of Metal-Carbon Bonds

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. 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.

79

Robertson, W. D.; Wang, M.; Warncke, K. J. Am. Chem. Soc. 2011, 133, 6968–6977. Jones, A. R.; Russell, H. J.; Greetham, G. M.; Towrie, M.; Hay, S.; Scrutton, N. S. J. Phys. Chem. A 2012, 116, 5586–5594. Russell, H. J.; Jones, A. R.; Hay, S.; Greetham, G. M.; Towrie, M.; Scrutton, N. S. Angew. Chem. Int. Ed. 2012, 51, 9306–9310. Jones, A. R.; Levy, C.; Hay, S.; Scrutton, N. S. FEBS J. 2013, 280, 2997–3008. Kutta, R. J.; Hardman, S. J. O.; Johannissen, L. O.; Bellina, B.; Messiha, H. L.; Manuel Ortiz-Guerrero, J.; Elias-Arnanz, M.; Padmanabhan, S.; Barran, P.; Scrutton, N. S.; Jones, A. R. Nat. Commun. 2015, 6, 1–11. Wiley, T. E.; Miller, W. R.; Miller, N. A.; Sension, R. J.; Lodowski, P.; Jaworska, M.; Kozlowski, P. M. J. Phys. Chem. Lett. 2016, 7, 143–147. Miller, N. A.; Wiley, T. E.; Spears, K. G.; Ruetz, M.; Kieninger, C.; Kraeutler, B.; Sension, R. J. J. Am. Chem. Soc. 2016, 138, 14250–14256. Miller, N. A.; Deb, A.; Alonso-Mori, R.; Glownia, J. M.; Kiefer, L. M.; Konar, A.; Michocki, L. B.; Sikorski, M.; Sofferman, D. L.; Song, S.; Toda, M. J.; Wiley, T. E.; Zhu, D.; Kozlowski, P. M.; Kubarych, K. J.; Penner-Hahn, J. E.; Sension, R. J. J. Phys. Chem. A 2018, 122, 4963–4971. Michocki, L. B.; Miller, N. A.; Alonso-Mori, R.; Britz, A.; Deb, A.; Glownia, J. M.; Kaneshiro, A. K.; Konar, A.; Koralek, J.; Meadows, J. H.; Sofferman, D. L.; Song, S.; Toda, M. J.; van Driel, T. B.; Kozlowski, P. M.; Kubarych, K. J.; Penner-Hahn, J. E.; Sension, R. J. J. Phys. Chem. B 2019, 123, 6042–6048. Lodowski, P.; Jaworska, M.; Andruniow, T.; Garabato, B. D.; Kozlowski, P. M. Phys. Chem. Chem. Phys. 2014, 16, 18675–18679. Lodowski, P.; Jaworska, M.; Andruniow, T.; Garabato, B. D.; Kozlowski, P. M. J. Phys. Chem. A 2014, 118, 11718–11734. Lodowski, P.; Jaworska, M.; Garabato, B. D.; Kozowski, P. M. J. Phys. Chem. A 2015, 119, 3913–3928. Garabato, B. D.; Kumar, N.; Lodowski, P.; Jaworska, M.; Kozlowski, P. M. Phys. Chem. Chem. Phys. 2016, 18, 4513–4526. Garabato, B. D.; Lodowski, P.; Jaworska, M.; Kozlowski, P. M. Phys. Chem. Chem. Phys. 2016, 18, 19070–19082. Kozlowski, P. M.; Garabato, B. D.; Lodowski, P.; Jaworska, M. Dalton Trans. 2016, 45, 4457–4470. Al Mamun, A.; Toda, M. J.; Lodowski, P.; Jaworska, M.; Kozlowski, P. M. ACS Catal. 2018, 8, 7164–7178. Ghosh, A. P.; Al Mamun, A.; Lodowski, P.; Jaworska, M.; Kozlowski, P. M. J. Photochem. Photobiol. B 2018, 189, 306–317. Al Mamun, A.; Toda, M. J.; Lodowski, P.; Kozlowski, P. M. J. Phys. Chem. B 2019, 123, 2585–2598. Al Mamun, A.; Toda, M. J.; Kozlowski, P. M. J. Photochem. Photobiol. B 2019, 191, 175–184. Toda, M. J.; Lodowski, P.; Al Mamun, A.; Jaworska, M.; Kozlowski, P. M. Coord. Chem. Rev. 2019, 385, 20–43. Hartung, J.; Hertel, B.; Trach, F. Chem. Ber. 1993, 126, 1187–1191. Fischer, H. Chem. Rev. 2001, 101, 3581–3610. Fischer, H. J. Am. Chem. Soc. 1986, 108, 3925–3927. Daikh, B. E.; Finke, R. G. J. Am. Chem. Soc. 1992, 114, 2938–2943. Maria, S.; Stoffelbach, F.; Mata, J.; Daran, J.-C.; Richard, P.; Poli, R. J. Am. Chem. Soc. 2005, 127, 5946–5956. Zhang, Y.; Schröder, K.; Kwak, Y.; Krys, P.; Morin, A. N.; Pintauer, T.; Poli, R.; Matyjaszewski, K. Macromolecules 2013, 46, 5512–5519. Tang, W.; Fukuda, T.; Matyjaszewski, K. Macromolecules 2006, 39, 4332–4337. Morales-Cerrada, R.; Fliedel, C.; Daran, J.-C.; Gayet, F.; Ladmiral, V.; Améduri, B.; Poli, R. Chem. Eur. J. 2019, 25, 296–308. Wayland, B. B.; Gridnev, A. A.; Ittel, S. D.; Fryd, M. Inorg. Chem. 1994, 33, 3830–3833. Woska, D. C.; Wayland, B. B. Inorg. Chim. Acta 1998, 270, 197–201. Geno, M. K.; Halpern, J. J. Am. Chem. Soc. 1987, 109, 1238–1240. Schrauzer, G. N.; Grate, J. H. J. Am. Chem. Soc. 1981, 103, 541–546. Kirker, G. W.; Bakac, A.; Espenson, J. H. J. Am. Chem. Soc. 1982, 104, 1249–1255. Schofield, M. H.; Halpern, J. Inorg. Chim. Acta 2003, 345, 353–358. Coombes, R. G.; Johnson, M. D. J. Chem. Soc. A 1966, 177–182. Schmidt, A. R.; Swaddle, T. W. J. Chem. Soc. A 1970, 1927–1932. Nohr, R. S.; Espenson, J. H. J. Am. Chem. Soc. 1975, 97, 3392–3396. Espenson, J. H.; Connolly, P.; Meyerstein, D.; Cohen, H. Inorg. Chem. 1983, 22, 1009–1013. Nappa, M. J.; Santi, R.; Diefenbach, S. P.; Halpern, J. J. Am. Chem. Soc. 1982, 104, 619–621. Nappa, M. J.; Santi, R.; Halpern, J. Organometallics 1985, 4, 34–41. Morales-Cerrada, R.; Fliedel, C.; Gayet, F.; Ladmiral, V.; Améduri, B.; Poli, R. Eur. J. Inorg. Chem. 2019, 4228–4233. Riordan, C. G.; Halpern, J. Inorg. Chim. Acta 1996, 243, 19–24. Collman, J. P.; McElweewhite, L.; Brothers, P. J.; Rose, E. J. Am. Chem. Soc. 1986, 108, 1332–1333. Halpern, J.; Ng, F. T. T.; Rempel, G. L. J. Am. Chem. Soc. 1979, 101, 7124–7126. Ng, F. T. T.; Rempel, G. L.; Mancuso, C.; Halpern, J. Organometallics 1990, 9, 2762–2772. Ng, F. T. T.; Rempel, G. L.; Halpern, J. J. Am. Chem. Soc. 1982, 104, 621–623. Ng, F. T. T.; Rempel, G. L.; Halpern, J. Inorg. Chim. Acta 1983, 77, L165–L166. Tsou, T.-T.; Loots, M.; Halpern, J. J. Am. Chem. Soc. 1982, 104, 623–624. Finke, R. G.; Smith, B. L.; Mayer, B. J.; Molinero, A. A. Inorg. Chem. 1983, 22, 3677–3679. Halpern, J.; Kim, S. H.; Leung, T. W. J. Am. Chem. Soc. 1984, 106, 8317–8319. Finke, R. G.; Hay, B. P. Inorg. Chem. 1984, 23, 3041–3043. Hay, B. P.; Finke, R. G. J. Am. Chem. Soc. 1986, 108, 4820–4829. Kim, S. H.; Chen, H. L.; Feilchenfeld, N.; Halpern, J. J. Am. Chem. Soc. 1988, 110, 3120–3126. Bakac, A.; Espenson, J. H. J. Am. Chem. Soc. 1984, 106, 5197–5202. Geno, M. K.; Halpern, J. J. Chem. Soc. Chem. Commun. 1987, 1052–1053. Schrauzer, G. N.; Holland, R. J. J. Am. Chem. Soc. 1971, 93, 1505–1506. Koenig, T.; Finke, R. G. J. Am. Chem. Soc. 1988, 110, 2657–2658. Marsh, E. N. G.; Ballou, D. P. Biochemistry 1998, 37, 11864–11872. Padmakumar, R.; Padmakumar, R.; Banerjee, R. Biochemistry 1997, 36, 3713–3718. Chowdhury, S.; Banerjee, R. J. Am. Chem. Soc. 2000, 122, 5417–5418. Poli, R. C. R. Chim. 2021, 24, 147–175. Champouret, Y.; MacLeod, K. C.; Smith, K. M.; Poli, R. Organometallics 2010, 29, 3125–3132. Poli, R.; Shaver, M. P. Inorg. Chem. 2014, 53, 7580–7590. Endicott, J. F.; Balakrishnan, K. P.; Wong, C. L. J. Am. Chem. Soc. 1980, 102, 5519–5526. McHatton, R. C.; Espenson, J. H.; Bakac, A. J. Am. Chem. Soc. 1982, 104, 3531–3533. Johnson, M. D. Acc. Chem. Res. 1983, 16, 343–349. Bertin, D.; Gigmes, D.; Marque, S. R. A.; Tordo, P. Macromolecules 2005, 38, 2638–2650. Hodgson, J. L.; Lin, C. Y.; Coote, M. L.; Marque, S. R. A.; Matyjaszewski, K. Macromolecules 2010, 43, 3728–3743. Lin, C. Y.; Marque, S. R. A.; Matyjaszewski, K.; Coote, M. L. Macromolecules 2011, 44, 7568–7583. Schuh, H. H.; Fischer, H. Helv. Chim. Acta 1978, 61, 2130–2164.

80

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. 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.

Reversible Homolysis of Metal-Carbon Bonds

Cohen, H.; Meyerstein, D. Inorg. Chem. 1974, 13, 2434–2443. Blau, R. J.; Espenson, J. H.; Bakac, A. Inorg. Chem. 1984, 23, 3526–3528. Mulac, W. A.; Meyerstein, D. J. Am. Chem. Soc. 1982, 104, 4124–4128. Elroi, H.; Meyerstein, D. J. Am. Chem. Soc. 1978, 100, 5540–5548. Roche, T. S.; Endicott, J. F. Inorg. Chem. 1974, 13, 1575–1580. Tait, A. M.; Hoffman, M. Z.; Hayon, E. Int. J. Radiat. Phys. Chem. 1976, 8, 691–696. Endicott, J. F.; Netzel, T. L. J. Am. Chem. Soc. 1979, 101, 4000–4002. Mok, C. Y.; Endicott, J. F. J. Am. Chem. Soc. 1978, 100, 123–129. Meyerstein, D.; Schwarz, H. A. J. Chem. Soc. Faraday Trans. 1988, 84, 2933–2949. Van Eldik, R.; Cohen, H.; Meyerstein, D. Angew. Chem. Int. Ed. 1991, 30, 1158–1160. Van Eldik, R.; Cohen, H.; Meyerstein, D. Inorg. Chem. 1994, 33, 1566–1568. Martinho Simões, J. A.; Beauchamp, J. L. Chem. Rev. 1990, 90, 629–688. Skinner, H. A. Experimental Thermochemistry; Interscience: New York, 1962; vol. II. Conner, J. A.; Skinner, H. A.; Virmani, Y. J. Chem. Soc. Faraday Trans. 1972, 68, 1754–1763. Connor, J. A.; Zafaranimoattar, M. T.; Bickerton, J.; Elsaied, N. I.; Suradi, S.; Carson, R.; Altakhin, G.; Skinner, H. A. Organometallics 1982, 1, 1166–1174. Stevens, A. E. Fundamental Studies of Reactive Intermediates in Organometallic Chemistry; California Institute of Technology, 1981. Folga, E.; Ziegler, T. J. Am. Chem. Soc. 1993, 115, 5169–5176. Poli, R.; Rahaman, S. M. W.; Ladmiral, V.; Améduri, B. J. Organomet. Chem. 2018, 864, 12–18. Marzilli, L. G.; Toscano, P. J.; Randaccio, L.; Bresciani-Pahor, N.; Calligaris, M. J. Am. Chem. Soc. 1979, 101, 6754–6756. Halpern, J. Science 1985, 227, 869. Woska, D. C.; Xie, Z. D.; Gridnev, A. A.; Ittel, S. D.; Fryd, M.; Wayland, B. B. J. Am. Chem. Soc. 1996, 118, 9102–9109. Tang, W.; Tsarevsky, N. V.; Matyjaszewski, K. J. Am. Chem. Soc. 2006, 128, 1598–1604. Fischer, H. Macromolecules 1997, 30, 5666–5672. Ohno, K.; Tsujii, Y.; Miyamoto, T.; Fukuda, T.; Goto, M.; Kobayashi, K.; Akaike, T. Macromolecules 1998, 31, 1064–1069. Fischer, H. J. Polym. Sci. A Polym. Chem. 1999, 37, 1885–1901. Zerk, T. J.; Bernhardt, P. V. Inorg. Chem. 2017, 56, 5784–5792. Zerk, T. J.; Gahan, L. R.; Krenske, E. H.; Bernhardt, P. V. Polym. Chem. 2019, 10, 1460–1470. Fantin, M.; Lorandi, F.; Ribelli, T. G.; Fliedel, C.; Thevenin, L.; Isse, A. A.; Poli, R.; Matyjaszewski, K. Macromolecules 2019, 52, 4079–4090. Matyjaszewski, K.; Tsarevsky, N. V. J. Am. Chem. Soc. 2014, 136, 6513–6533. Boys, S. F.; Bernardi, F. Mol. Phys. 1970, 19, 553–566. Davidson, E. R.; Feller, D. Chem. Rev. 1986, 86, 681–696. Salomon, O.; Reiher, M.; Hess, B. A. J. Chem. Phys. 2002, 117, 4729–4737. Andruniow, T.; Zgierski, M. Z.; Kozlowski, P. M. J. Am. Chem. Soc. 2001, 123, 2679–2680. Dölker, N.; Maseras, F.; Lledós, A. J. Phys. Chem. B 2001, 105, 7564–7571. Jensen, K. P.; Ryde, U. J. Phys. Chem. A 2003, 107, 7539–7545. Siegbahn, P. E. M.; Blomberg, M. R. A.; Chen, S. L. J. Chem. Theory Comput. 2010, 6, 2040–2044. Kuta, J.; Patchkovskii, S.; Zgierski, M. Z.; Kozlowski, P. M. J. Comput. Chem. 2006, 27, 1429–1437. Chen, S. L.; Blomberg, M. R. A.; Siegbahn, P. E. M. J. Phys. Chem. B 2011, 115, 4066–4077. Jensen, K. P.; Roos, B. O.; Ryde, U. J. Chem. Phys. 2007, 126, 014103-1–014103-14. Kozlowski, P. M.; Kumar, M.; Piecuch, P.; Li, W.; Bauman, N. P.; Hansen, J. A.; Lodowski, P.; Jaworska, M. J. Chem. Theory Comput. 2012, 8, 1870–1894. Morin, A. N.; Detrembleur, C.; Jérôme, C.; Tullio, P. D.; Poli, R.; Debuigne, A. Macromolecules 2013, 46, 4303–4312. Qi, X. J.; Li, Z.; Fu, Y.; Guo, Q. X.; Liu, L. Organometallics 2008, 27, 2688–2698. Ragsdale, S. W. Chem. Rev. 2006, 106, 3317–3337. Bridwell-Rabb, J.; Grell, T. A. J.; Drennan, C. L. Annu. Rev. Biochem. 2018, 87, 555–584. Broderick, W. E.; Broderick, J. B. J. Biol. Inorg. Chem. 2019, 24, 769–776. Broderick, W. E.; Hoffman, B. M.; Broderick, J. B. Acc. Chem. Res. 2018, 51, 2611–2619. Matthews, R. G. Acc. Chem. Res. 2001, 34, 681–689. Thauer, R. K. Biochemistry 2019, 58, 5198–5220. Can, M.; Armstrong, F. A.; Ragsdale, S. W. Chem. Rev. 2014, 114, 4149–4174. Obeid, R. Vitamin B12: Advances and Insights, 1st ed.; CRC Press: Boca Raton, FL, 2017; p 376. Gruber, K.; Puffer, B.; Kräutler, B. Chem. Soc. Rev. 2011, 40, 4346–4363. Brown, K. L. Chem. Rev. 2005, 105, 2075–2150. Butler, P. A.; Kräutler, B. Top. Organomet. Chem. 2006, 17, 1–55. Marsh, E. N. G.; Drennan, C. L. Curr. Opin. Chem. Biol. 2001, 5, 499–505. Toraya, T. Chem. Rev. 2003, 103, 2095–2128. Sandala, G. M.; Smith, D. M.; Radom, L. Acc. Chem. Res. 2010, 43, 642–651. Wu, B.; Szymanski, W.; Heberling, M. M.; Feringa, B. L.; Janssen, D. B. Trends Biotechnol. 2011, 29, 352–362. Hodgkin, D. C.; Kamper, J.; Mackay, M.; Pickworth, J.; Trueblood, K. N.; White, J. G. Nature 1956, 178, 64–66. Lenhert, P. G.; Hodgkin, D. C. Nature 1961, 192, 937–938. Rossi, M.; Glusker, J. P.; Randaccio, L.; Summers, M. F.; Toscano, P. J.; Marzilli, L. G. J. Am. Chem. Soc. 1985, 107, 1729–1738. Woodward, R. B. Pure Appl. Chem. 1973, 33, 145–177. Eschenmoser, A.; Wintner, C. E. Science 1977, 196, 1410–1420. Beaven, G. H.; Johnson, E. A. Nature 1955, 176, 1264–1265. Lester Smith, E.; Mervyn, L.; Johnson, A. W.; Shaw, N. Nature 1962, 194, 1175. Schrauzer, G. N.; Hashimoto, T. J. Am. Chem. Soc. 1979, 101, 4593–4601. Ruetz, M.; Gherasim, C.; Gruber, K.; Fedosov, S.; Banerjee, R.; Kräutler, B. Angew. Chem. Int. Ed. 2013, 52, 2606–2610. Ruetz, M.; Salchner, R.; Wurst, K.; Fedosov, S.; Kräutler, B. Angew. Chem. Int. Ed. 2013, 52, 11406–11409. Giedyk, M.; Goliszewska, K.; Gryko, D. Chem. Soc. Rev. 2015, 44, 3391–3404. Banerjee, R.; Ragsdale, S. W. Annu. Rev. Biochem. 2003, 72, 209–247. Jensen, K. P.; Ryde, U. J. Am. Chem. Soc. 2005, 127, 9117–9128. Conrad, K. S.; Jordan, C. D.; Brown, K. L.; Brunold, T. C. Inorg. Chem. 2015, 54, 3736–3747. Banerjee, R. Chem. Rev. 2003, 103, 2083–2094. Frey, P. A.; Essenberg, M. K.; Abeles, R. H. J. Biol. Chem. 1967, 242, 5369–5377.

Reversible Homolysis of Metal-Carbon Bonds

176. 177. 178. 179. 180. 181. 182. 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. 233. 234. 235. 236. 237. 238. 239. 240. 241. 242. 243. 244.

81

Hay, B. P.; Finke, R. G. J. Am. Chem. Soc. 1987, 109, 8012–8018. Wolthers, K. R.; Rigby, S. E. J.; Scrutton, N. S. J. Biol. Chem. 2008, 283, 34615–34625. Chowdhury, S.; Banerjee, R. Biochemistry 2000, 39, 7998–8006. Abeles, R. H.; Dolphin, D. Acc. Chem. Res. 1976, 9, 114–120. Barker, H. A.; Weissbach, H.; Smyth, R. D. Proc. Natl. Acad. Sci. U. S. A. 1958, 44, 1093–1097. Bucher, D.; Sandala, G. M.; Durbeej, B.; Radom, L.; Smith, D. M. J. Am. Chem. Soc. 2012, 134, 1591–1599. Kumar, N.; Bucher, D.; Kozlowski, P. M. J. Phys. Chem. B 2019, 123, 2210–2216. Ruetz, M.; Campanello, G. C.; Purchal, M.; Shen, H.; McDevitt, L.; Gouda, H.; Wakabayashi, S.; Zhu, J.; Rubin, E. J.; Warncke, K.; Mootha, V. K.; Koutmos, M.; Banerjee, R. Science 2019, 366, 589–593. Stadtman, T. C. In The Enzymes, 3rd ed.; Boyer, P. D., Ed.; Academic Press: New York and London, 1972; vol. 6; pp 539–563. Frey, P. A. Acc. Chem. Res. 2014, 47, 540–549. Toraya, T. Cell. Mol. Life Sci. 2000, 57, 106–127. Finlay, T. H.; Valinsky, J.; Mildvan, A. S.; Abeles, R. H. J. Biol. Chem. 1973, 248, 1285–1290. Yamanishi, M.; Ide, H.; Murakami, Y.; Toraya, T. Biochemistry 2005, 44, 2113–2118. Kamachi, T.; Toraya, T.; Yoshizawa, K. J. Am. Chem. Soc. 2004, 126, 16207–16216. Toraya, T.; Tanokuchi, A.; Yamasaki, A.; Nakamura, T.; Ogura, K.; Tobimatsu, T. Biochemistry 2016, 55, 69–78. Abend, A.; Nitsche, R.; Bandarian, V.; Stupperich, E.; Rétey, J. Angew. Chem. Int. Ed. 1998, 37, 625–627. Abend, A.; Bandarian, V.; Nitsche, R.; Stupperich, E.; Rétey, J.; Reed, G. H. Arch. Biochem. Biophys. 1999, 370, 138–141. Zhu, C.; Warncke, K. J. Am. Chem. Soc. 2010, 132, 9610–9615. Kohne, M.; Li, W.; Zhu, C.; Warncke, K. Biochemistry 2019, 58, 3683–3690. Wetmore, S. D.; Smith, D. M.; Bennett, J. T.; Radom, L. J. Am. Chem. Soc. 2002, 124, 14054–14065. Semialjac, M.; Schwarz, H. J. Am. Chem. Soc. 2002, 124, 8974–8983. Semialjac, M.; Schwarz, H. J. Org. Chem. 2003, 68, 6967–6983. Semialjac, M.; Schwarz, H. Chem. Eur. J. 2004, 10, 2781–2788. Shibata, N.; Tamagaki, H.; Hieda, N.; Akita, K.; Komori, H.; Shomura, Y.; Terawaki, S.-I.; Mori, K.; Yasuoka, N.; Higuchi, Y.; Toraya, T. J. Biol. Chem. 2010, 285, 26484–26493. Mori, K.; Oiwa, T.; Kawaguchi, S.; Kondo, K.; Takahashi, Y.; Toraya, T. Biochemistry 2014, 53, 2661–2671. Wolthers, K. R.; Levy, C.; Scrutton, N. S.; Leys, D. J. Biol. Chem. 2010, 285, 13942–13950. Berkovitch, F.; Behshad, E.; Tang, K.-H.; Enns, E. A.; Frey, P. A.; Drennan, C. L. Proc. Natl. Acad. Sci. 2004, 101, 15870–15875. Tang, K.-H.; Casarez, A. D.; Wu, W.; Frey, P. A. Arch. Biochem. Biophys. 2003, 418, 49–54. Wetmore, S. D.; Smith, D. M.; Radom, L. J. Am. Chem. Soc. 2001, 123, 8678–8689. Chen, Y.-H.; Maity, A. N.; Frey, P. A.; Ke, S.-C. J. Am. Chem. Soc. 2013, 135, 788–794. Lo, H.-H.; Lin, H.-H.; Maity, A. N.; Ke, S.-C. Chem. Commun. 2016, 52, 6399–6402. Sandala, G. M.; Smith, D. M.; Radom, L. J. Am. Chem. Soc. 2006, 128, 16004–16005. Chen, Y.-H.; Maity, A. N.; Pan, Y.-C.; Frey, P. A.; Ke, S.-C. J. Am. Chem. Soc. 2011, 133, 17152–17155. Stubbe, J.; van der Donk, W. A. Chem. Rev. 1998, 98, 705–762. Mulliez, E.; Fontecave, M. Coord. Chem. Rev. 1999, 185–186, 775–793. Kolberg, M.; Strand, K. R.; Graff, P.; Kristoffer Andersson, K. Biochimica et Biophysica Acta, Proteins and Proteomics 2004, 1699, 1–34. Lawrence, C. C.; Stubbe, J. Curr. Opin. Chem. Biol. 1998, 2, 650–655. Booker, S.; Licht, S.; Broderick, J.; Stubbe, J. Biochemistry 1994, 33, 12676–12685. Licht, S.; Gerfen, G. J.; Stubbe, J. Science 1996, 271, 477–481. Gerfen, G. J.; Licht, S.; Willems, J.-P.; Hoffman, B. M.; Stubbe, J. J. Am. Chem. Soc. 1996, 118, 8192–8197. Sintchak, M. D.; Arjara, G.; Kellogg, B. A.; Stubbe, J.; Drennan, C. L. Nat. Struct. Biol. 2002, 9, 293–300. Larsson, K.-M.; Logan, D. T.; Nordlund, P. ACS Chem. Biol. 2010, 5, 933–942. Robins, M. J.; Ewing, G. J. J. Am. Chem. Soc. 1999, 121, 5823–5824. Lenz, R.; Giese, B. J. Am. Chem. Soc. 1997, 119, 2784–2794. Broderick, J. B.; Duffus, B. R.; Duschene, K. S.; Shepard, E. M. Chem. Rev. 2014, 114, 4229–4317. Horitani, M.; Shisler, K.; Broderick, W. E.; Hutcheson, R. U.; Duschene, K. S.; Marts, A. R.; Hoffman, B. M.; Broderick, J. B. Science 2016, 352, 822–825. Byer, A. S.; Yang, H.; McDaniel, E. C.; Kathiresan, V.; Impano, S.; Pagnier, A.; Watts, H.; Denler, C.; Vagstad, A. L.; Piel, J.; Duschene, K. S.; Shepard, E. M.; Shields, T. P.; Scott, L. G.; Lilla, E. A.; Yokoyama, K.; Broderick, W. E.; Hoffman, B. M.; Broderick, J. B. J. Am. Chem. Soc. 2018, 140, 8634–8638. Yang, H.; McDaniel, E. C.; Impano, S.; Byer, A. S.; Jodts, R. J.; Yokoyama, K.; Broderick, W. E.; Broderick, J. B.; Hoffman, B. M. J. Am. Chem. Soc. 2019, 141, 12139–12146. Ye, M.; Thompson, N. B.; Brown, A. C.; Suess, D. L. M. J. Am. Chem. Soc. 2019, 141, 13330–13335. Brown, A. C.; Suess, D. L. M. J. Am. Chem. Soc. 2020, 142, 14240–14248. Dong, M.; Kathiresan, V.; Fenwick, M. K.; Torelli, A. T.; Zhang, Y.; Caranto, J. D.; Dzikovski, B.; Sharma, A.; Lancaster, K. M.; Freed, J. H.; Ealick, S. E.; Hoffman, B. M.; Lin, H. Science 2018, 359, 1247. Ram, M. S.; Riordan, C. G. J. Am. Chem. Soc. 1995, 117, 2365–2366. Ram, M. S.; Riordan, C. G.; Yap, G. P. A.; Liable-Sands, L.; Rheingold, A. L.; Marchaj, A.; Norton, J. R. J. Am. Chem. Soc. 1997, 119, 1648–1655. Harder, S. R.; Lu, W. P.; Feinberg, B. A.; Ragsdale, S. W. Biochemistry 1989, 28, 9080–9087. Martin, B. D.; Finke, R. G. J. Am. Chem. Soc. 1990, 112, 2419–2420. Jaun, B. Helv. Chim. Acta 1990, 73, 2209–2217. Berkessel, A. Bioorg. Chem. 1991, 19, 101–115. Wongnate, T.; Sliwa, D.; Ginovska, B.; Smith, D.; Wolf, M. W.; Lehnert, N.; Raugei, S.; Ragsdale, S. W. Science 2016, 352, 953. Yan, M.; Lo, J. C.; Edwards, J. T.; Baran, P. S. J. Am. Chem. Soc. 2016, 138, 12692–12714. Iqbal, J.; Bhatia, B.; Nayyar, N. K. Chem. Rev. 1994, 94, 519–564. Bolm, C.; Legros, J.; Le Paih, J.; Zani, L. Chem. Rev. 2004, 104, 6217–6254. Bauer, I.; Knölker, H.-J. Chem. Rev. 2015, 115, 3170–3387. Gualandi, A.; Mengozzi, L.; Cozzi, P. G. Asian J. Org. Chem. 2017, 6, 1160–1179. Jahn, U. In Radicals in Synthesis III; Heinrich, M. R., Gansauer, A., Eds.; Springer, 2012; vol. 320; pp 191–322. Crossley, S. W. M.; Obradors, C.; Martinez, R. M.; Shenvi, R. A. Chem. Rev. 2016, 116, 8912–9000. Kyne, S. H.; Lefèvre, G.; Ollivier, C.; Petit, M.; Ramis Cladera, V.-A.; Fensterbank, L. Chem. Soc. Rev. 2020, 49, 8501–8542. Pattenden, G. Chem. Soc. Rev. 1988, 17, 361–382. te Grotenhuis, C.; de Bruin, B. Synlett 2018, 29, 2238–2250. Rossi, B.; Prosperini, S.; Pastori, N.; Clerici, A.; Punta, C. Molecules 2012, 17, 14700–14732.

82

Reversible Homolysis of Metal-Carbon Bonds

245. Rosales, A.; Rodriguez-Garcia, I.; Munoz-Bascon, J.; Roldan-Molina, E.; Padial, N. M.; Pozo Morales, L.; Garcia-Ocana, M.; Enrique Oltra, J. Eur. J. Org. Chem. 2015, 2015, 4567–4591. 246. Manuel Botubol-Ares, J.; Jesus Duran-Pena, M.; Hanson, J. R.; Hernandez-Galan, R.; Collado, I. G. Synthesis-Stuttgart 2018, 50, 2163–2180. 247. Wessjohann, L. A.; Scheid, G. Synthesis-Stuttgart 1999, 1–36. 248. Melikyan, G. G. Synthesis-Stuttgart 1993, 833–850. 249. Demir, A. S.; Emrullahoglu, M. Curr. Org. Synth. 2007, 4, 323–351. 250. Pan, X.-Q.; Zou, J.-P.; Zhang, W. Mol. Divers. 2009, 13, 421–438. 251. Mondal, M.; Bora, U. RSC Adv. 2013, 3, 18716–18754. 252. Cheng, L.-J.; Mankad, N. P. Chem. Soc. Rev. 2020, 49, 8036–8064. 253. Studer, A.; Curran, D. P. Angew. Chem. Int. Ed. 2016, 55, 58–102. 254. Feder, H. M.; Halpern, J. J. Am. Chem. Soc. 1975, 97, 7186–7188. 255. Sweany, R. L.; Halpern, J. J. Am. Chem. Soc. 1977, 99, 8335–8337. 256. Samsel, E. G.; Kochi, J. K. J. Am. Chem. Soc. 1986, 108, 4790–4804. 257. Munoz-Molina, J. M.; Belderrain, T. R.; Perez, P. J. Eur. J. Inorg. Chem. 2011, 3155–3164. 258. Halpern, J. Acc. Chem. Res. 1970, 3, 386–392. 259. Davis, K. A.; Matyjaszewski, K. J. Macromol. Sci. A 2004, A41, 449–465. 260. Rajanbabu, T. V.; Nugent, W. A. J. Am. Chem. Soc. 1994, 116, 986–997. 261. Hilt, G.; Bolze, P.; Kieltsch, I. Chem. Commun. 2005, 1996–1998. 262. Hilt, G.; Bolze, P.; Harms, K. Chem. Eur. J. 2007, 13, 4312–4325. 263. Okabe, M.; Tada, M. Chem. Lett. 1980, 831–834. 264. Tada, M.; Okabe, M. Chem. Lett. 1980, 201–204. 265. Okabe, M.; Abe, M.; Tada, M. J. Org. Chem. 1982, 47, 1775–1777. 266. Okabe, M.; Tada, M. J. Org. Chem. 1982, 47, 5382–5384. 267. Torii, S.; Inokuchi, T.; Yukawa, T. J. Org. Chem. 1985, 50, 5875–5877. 268. Bhandal, H.; Pattenden, G.; Russell, J. J. Tetrahedron Lett. 1986, 27, 2299–2302. 269. Patel, V. F.; Pattenden, G.; Russell, J. J. Tetrahedron Lett. 1986, 27, 2303–2306. 270. Begley, M. J.; Bhandal, H.; Hutchinson, J. H.; Pattenden, G. Tetrahedron Lett. 1987, 28, 1317–1320. 271. Baldwin, J. E.; Li, C. S. J. Chem. Soc. Chem. Commun. 1987, 166–168. 272. Branchaud, B. P.; Detlefsen, W. D. Tetrahedron Lett. 1991, 32, 6273–6276. 273. Giese, B.; Erdmann, P.; Gobel, T.; Springer, R. Tetrahedron Lett. 1992, 33, 4545–4548. 274. Schrauzer, G. N.; Deutsch, E. J. Am. Chem. Soc. 1969, 91, 3341–3350. 275. Zhou, D. L.; Walder, P.; Scheffold, R.; Walder, L. Helv. Chim. Acta 1992, 75, 995–1011. 276. Felkin, H.; Knowles, P. J.; Meunier, B. J. Organomet. Chem. 1978, 146, 151–167. 277. Ekomié, A.; Lefèvre, G.; Fensterbank, L.; Lacôte, E.; Malacria, M.; Ollivier, C.; Jutand, A. Angew. Chem. Int. Ed. 2012, 51, 6942–6946. 278. Kyne, S. H.; Clémancey, M.; Blondin, G.; Derat, E.; Fensterbank, L.; Jutand, A.; Lefèvre, G.; Ollivier, C. Organometallics 2018, 37, 761–771. 279. de Bruin, B.; Hetterscheid, D. G. H. Eur. J. Inorg. Chem. 2007, 211–230. 280. Shevick, S. L.; Wilson, C. V.; Kotesova, S.; Kim, D.; Holland, P. L.; Shenvi, R. A. Chem. Sci. 2020, 11, 12401–12422. 281. Jiang, H.; Lai, W.; Chen, H. ACS Catal. 2019, 9, 6080–6086. 282. Eisenberg, D. C.; Norton, J. R. Isr. J. Chem. 1991, 31, 55–66. 283. Wang, D.; Angelici, R. J. J. Am. Chem. Soc. 1996, 118, 935–942. 284. Lo, J. C.; Yabe, Y.; Baran, P. S. J. Am. Chem. Soc. 2014, 136, 1304–1307. 285. Lo, J. C.; Kim, D.; Pan, C. M.; Edwards, J. T.; Yabe, Y.; Gui, J. H.; Qin, T.; Gutierrez, S.; Giacoboni, J.; Smith, M. W.; Holland, P. L.; Baran, P. S. J. Am. Chem. Soc. 2017, 139, 2484–2503. 286. Kim, D.; Rahaman, S. M. W.; Mercado, B. Q.; Poli, R.; Holland, P. L. J. Am. Chem. Soc. 2019, 141, 7473–7485. 287. Branchaud, B. P.; Yu, G. X. Organometallics 1993, 12, 4262–4264. 288. Patel, V. F.; Pattenden, G. J. Chem. Soc. Chem. Commun. 1987, 871–872. 289. Patel, V. F.; Pattenden, G. Tetrahedron Lett. 1987, 28, 1451–1454. 290. Ghosez, A.; Gobel, T.; Giese, B. Chem. Ber. 1988, 121, 1807–1811. 291. Branchaud, B. P.; Meier, M. S.; Choi, Y. Tetrahedron Lett. 1988, 29, 167–170. 292. Branchaud, B. P.; Choi, Y. L. J. Org. Chem. 1988, 53, 4638–4641. 293. Branchaud, B. P.; Yu, G. X. Tetrahedron Lett. 1988, 29, 6545–6548. 294. Bhandal, H.; Howell, A. R.; Patel, V. F.; Pattenden, G. J. Chem. Soc. Perkin Trans. 1990, 1, 2709–2714. 295. Ali, A.; Harrowven, D. C.; Pattenden, G. Tetrahedron Lett. 1992, 33, 2851–2854. 296. Coveney, D. J.; Patel, V. F.; Pattenden, G. Tetrahedron Lett. 1987, 28, 5949–5952. 297. Patel, V. F.; Pattenden, G. Tetrahedron Lett. 1988, 29, 707–710. 298. Gill, G. B.; Pattenden, G.; Reynolds, S. J. Tetrahedron Lett. 1989, 30, 3229–3232. 299. Gill, G. B.; Pattenden, G.; Reynolds, S. J. J. Chem. Soc. Perkin Trans. 1994, 1, 369–378. 300. Kwiatek, J.; Seyler, J. K. Adv. Chem. Ser. 1968, (70), 207–232. 301. Halpern, J.; Wong, L.-Y. J. Am. Chem. Soc. 1968, 90, 6665–6669. 302. Chung, S. K. J. Org. Chem. 1979, 44, 1014–1016. 303. Magnus, P.; Waring, M. J.; Scott, D. A. Tetrahedron Lett. 2000, 41, 9731–9733. 304. Iwasaki, K.; Wan, K. K.; Oppedisano, A.; Crossley, S. W. M.; Shenvi, R. A. J. Am. Chem. Soc. 2014, 136, 1300–1303. 305. King, S. M.; Ma, X.; Herzon, S. B. J. Am. Chem. Soc. 2014, 136, 6884–6887. 306. Crossley, S. W. M.; Barabé, F.; Shenvi, R. A. J. Am. Chem. Soc. 2014, 136, 16788–16791. 307. Kräutler, B. Helv. Chim. Acta 1984, 67, 1053–1059. 308. Gridnev, A. A.; Ittel, S. D. Chem. Rev. 2001, 101, 3611–3659. 309. Hatakeyama, T.; Hashimoto, T.; Kondo, Y.; Fujiwara, Y.; Seike, H.; Takaya, H.; Tamada, Y.; Ono, T.; Nakamura, M. J. Am. Chem. Soc. 2010, 132, 10674–10676. 310. Daifuku, S. L.; Al-Afyouni, M. H.; Snyder, B. E. R.; Kneebone, J. L.; Neidig, M. L. J. Am. Chem. Soc. 2014, 136, 9132–9143. 311. Przyojski, J. A.; Veggeberg, K. P.; Arman, H. D.; Tonzetich, Z. J. ACS Catal. 2015, 5, 5938–5946. 312. Liu, Y. S.; Xiao, J.; Wang, L.; Song, Y.; Deng, L. Organometallics 2015, 34, 599–605. 313. Thevenin, L.; Fliedel, C.; Matyjaszewski, K.; Poli, R. Eur. J. Inorg. Chem. 2019, 4489–4499. 314. Parchomyk, T.; Koszinowski, K. Synthesis-Stuttgart 2017, 49, 3269–3280. 315. Cahiez, G.; Moyeux, A. Chem. Rev. 2010, 110, 1435–1462.

Reversible Homolysis of Metal-Carbon Bonds

316. 317. 318. 319. 320. 321. 322. 323. 324. 325. 326. 327. 328. 329. 330. 331. 332. 333. 334. 335. 336. 337. 338. 339. 340. 341. 342. 343. 344. 345. 346. 347. 348. 349. 350. 351. 352. 353. 354. 355. 356. 357. 358. 359. 360. 361. 362. 363. 364. 365. 366. 367. 368. 369. 370. 371. 372. 373. 374. 375. 376. 377. 378. 379. 380. 381. 382. 383. 384. 385. 386. 387. 388.

Diccianni, J. B.; Diao, T. Trends Chem. 2019, 1, 830–844. Sambiagio, C.; Marsden, S. P.; Blacker, A. J.; McGowan, P. C. Chem. Soc. Rev. 2014, 43, 3525–3550. Kleimark, J.; Hedstrom, A.; Larsson, P.-F.; Johansson, C.; Norrby, P.-O. ChemCatChem 2009, 1, 152–161. Lefèvre, G.; Jutand, A. Chem. Eur. J. 2014, 20, 4796–4805. Nakao, J.; Inoue, R.; Shinokubo, H.; Oshima, K. J. Org. Chem. 1997, 62, 1910–1911. Howell, A. R.; Pattenden, G. J. Chem. Soc., Perkin Trans. 1 1990, 2715–2719. Howell, A. R.; Pattenden, G. J. Chem. Soc. Chem. Commun. 1990, 103–104. Tokuyasu, T.; Kunikawa, S.; Masuyama, A.; Nojima, M. Org. Lett. 2002, 4, 3595–3598. Sugimori, T.; Horike, S.; Tsumura, S.; Handa, M.; Kasuga, K. Inorg. Chim. Acta 1998, 283, 275–278. Kendrick, M. J.; Al-Akhdar, W. Inorg. Chem. 1987, 26, 3971–3972. Mikolaiski, W.; Baum, G.; Massa, W.; Hoffmann, R. W. J. Organomet. Chem. 1989, 376, 397–405. Bhuyan, M.; Laskar, M.; Mandal, D.; Gupta, B. D. Organometallics 2007, 26, 3559–3567. Zhao, Y.; Wang, Y.; Zhou, X.; Xue, Z.; Wang, X.; Xie, X.; Poli, R. Angew. Chem. Int. Ed. 2019, 58, 14311–14318. Waser, J.; Carreira, E. M. J. Am. Chem. Soc. 2004, 126, 5676–5677. Waser, J.; Gaspar, B.; Nambu, H.; Carreira, E. M. J. Am. Chem. Soc. 2006, 128, 11693–11712. Shigehisa, H.; Aoki, T.; Yamaguchi, S.; Shimizu, N.; Hiroya, K. J. Am. Chem. Soc. 2013, 135, 10306–10309. Shigehisa, H.; Koseki, N.; Shimizu, N.; Fujisawa, M.; Niitsu, M.; Hiroya, K. J. Am. Chem. Soc. 2014, 136, 13534–13537. Shigehisa, H.; Nishi, E.; Fujisawa, M.; Hiroya, K. Org. Lett. 2013, 15, 5158–5161. Zhou, X.-L.; Yang, F.; Sun, H.-L.; Yin, Y.-N.; Ye, W.-T.; Zhu, R. J. Am. Chem. Soc. 2019, 141, 7250–7255. Gaspar, B.; Carreira, E. M. Angew. Chem. Int. Ed. 2007, 46, 4519–4522. Gaspar, B.; Carreira, E. M. Angew. Chem. Int. Ed. 2008, 47, 5758–5760. Gaspar, B.; Carreira, E. M. J. Am. Chem. Soc. 2009, 131, 13214–13215. Girijavallabhan, V.; Alvarez, C.; Njoroge, F. G. J. Org. Chem. 2011, 76, 6442–6446. Matos, J. L. M.; Vasquez-Cespedes, S.; Gu, J.; Oguma, T.; Shenvi, R. A. J. Am. Chem. Soc. 2018, 140, 16976–16981. Scheffold, R.; Abrecht, S.; Orlinski, R.; Ruf, H. R.; Stamouli, P.; Tinembart, O.; Walder, L.; Weymuth, C. Pure Appl. Chem. 1987, 59, 363–372. Jenkins, A. D.; Jones, R. G.; Moad, G. Pure Appl. Chem. 2010, 82, 483–491. Hawker, C. J.; Bosman, A. W.; Harth, E. Chem. Rev. 2001, 101, 3661–3688. David, G.; Boyer, C.; Tonnar, J.; Améduri, B.; Lacroix-Desmazes, P.; Boutevin, B. Chem. Rev. 2006, 106, 3936–3962. Poli, R. Angew. Chem. Int. Ed. 2006, 45, 5058–5070. Poli, R. In Polymer Science: A Comprehensive Reference; Matyjaszewski, K., Möller, M., Eds.; Elsevier BV: Amsterdam, 2012; vol. 3; pp 351–375. Poli, R. Reference Module in Materials Science and Materials Engineering; Elsevier, 2016. Matyjaszewski, K. Adv. Mater. 2018, 30, 1706441. Ouchi, M.; Sawamoto, M. Macromolecules 2017, 50, 2603–2614. Debuigne, A.; Poli, R.; Jérôme, C.; Jérome, R.; Detrembleur, C. Prog. Polym. Sci. 2009, 34, 211–239. Smith, K. M.; McNeil, W. S.; Abd-El-Aziz, A. S. Macromol. Chem. Phys. 2010, 211, 10–16. Hurtgen, M.; Detrembleur, C.; Jerome, C.; Debuigne, A. Polym. Rev. 2011, 51, 188–213. Allan, L. E. N.; Perry, M. R.; Shaver, M. P. Prog. Polym. Sci. 2012, 37, 127–156. Peng, C.-H.; Yang, T.-Y.; Zhao, Y.; Fu, X. Org. Biomol. Chem. 2014, 12, 8580–8587. Poli, R.; Allan, L. E. N.; Shaver, M. P. Prog. Polym. Sci. 2014, 39, 1827–1845. Poli, R. Chem. Eur. J. 2015, 21, 6988–7001. Debuigne, A.; Jerome, C.; Detrembleur, C. Polymer 2017, 115, 285–307. Fliedel, C.; Poli, R. J. Organomet. Chem. 2019, 880, 241–252. Odian, G. Principles of Polymerization, 4th ed.; John Wiley & Sons, Inc.: Hoboken, NJ, 2004. Asandei, A. D.; Moran, I. W. J. Am. Chem. Soc. 2004, 126, 15932–15933. Asandei, A. D.; Moran, I. W.; Castro, M. A. Polym. Prepr. 2003, 44, 829–830. Asandei, A. D.; Chen, Y. Polym. Prepr. 2004, 45, 766–767. Asandei, A. D.; Moran, I. W. J. Polym. Sci. Polym. Chem. 2005, 43, 6039–6047. Asandei, A. D.; Moran, I. W. J. Polym. Sci. Polym. Chem. 2006, 44, 1060–1070. Asandei, A. D.; Moran, I. W.; Saha, G.; Chen, Y. J. Polym. Sci. Polym. Chem. 2006, 44, 2156–2165. Asandei, A. D.; Saha, G. PMSE Prepr. 2006, 94, 480–481. Asandei, A. D.; Saha, G. J. Polym. Sci. A Polym. Chem. 2006, 44, 1106–1116. Asandei, A. D.; Moran, I. W.; Saha, G.; Chen, Y. ACS Symp. Ser. 2006, 944, 125–139. Asandei, A. D.; Saha, G. Polym. Prepr. 2007, 48, 272–273. Asandei, A. D.; Chen, Y. PMSE Prepr. 2007, 97, 450–451. Asandei, A. D.; Chen, Y.; Moran, I. W.; Saha, G. J. Organomet. Chem. 2007, 692, 3174–3182. Asandei, A. D.; Chen, Y.; Simpson, C.; Gilbert, M.; Moran, I. W. Polym. Prepr. 2008, 49, 489–490. Asandei, A. D.; Chen, Y.; Saha, G.; Moran, I. W. Tetrahedron 2008, 64, 11831–11838. Asandei, A. D.; Chen, Y.; Adebolu, O. PMSE Prepr. 2008, 98, 370–371. Asandei, A. D.; Yu, H. S.; Adebolu, O. PMSE Prepr. 2009, 101, 1377–1378. Asandei, A. D.; Yu, H. S.; Simpson, C. P. PMSE Prepr. 2009, 101, 1379–1380. Asandei, A. D.; Yu, H. S. Polym. Prepr. 2009, 50, 450–451. Asandei, A. D.; Yu, H. S.; Simpson, C. P. PMSE Prepr. 2010, 103, 511–512. Asandei, A. D.; Yu, H. S.; Simpson, C. P. PMSE Prepr. 2010, 102, 68–69. Asandei, A. D.; Simpson, C. P.; Yu, H. S. Polym. Prepr. 2008, 49, 73–74. Asandei, A. D.; Simpson, C. P.; Yu, H. S.; Adebolu, O. I.; Saha, G.; Chen, Y. ACS Symp. Ser. 2009, 1024, 149–163. Asandei, A. D.; Moran, I. W. J. Polym. Sci. A Polym. Chem. 2005, 43, 6028–6038. Asandei, A. D.; Moran, I. W.; Saha, G.; Chen, Y. H. J. Polym. Sci. A Polym. Chem. 2006, 44, 2015–2026. Asandei, A. D.; Chen, Y. Polym. Prepr. 2007, 48, 232–233. Asandei, A.; Saha, G. Polym. Mater. Sci. Eng. 2005, 93, 470–471. Asandei, A. D.; Moran, I. W. Polym. Prepr. 2005, 46, 136–137. Coward, D. L.; Lake, B. R. M.; Poli, R.; Shaver, M. P. Macromolecules 2019, 52, 3252–3256. Grishin, D.; Semyonycheva, L.; Telegina, E.; Smirnov, A.; Nevodchikov, V. Russ. Chem. Bull. 2003, 52, 505–507. Grishin, D. F.; Ignatov, S. K.; Shchepalov, A. A.; Razuvaev, A. G. Appl. Organomet. Chem. 2004, 18, 271–276.

83

84

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. 414. 415. 416. 417. 418. 419. 420. 421. 422. 423. 424. 425. 426. 427. 428. 429. 430. 431. 432. 433. 434. 435. 436. 437. 438. 439. 440. 441. 442. 443. 444. 445. 446. 447. 448. 449. 450. 451. 452. 453. 454. 455. 456. 457. 458. 459. 460. 461.

Reversible Homolysis of Metal-Carbon Bonds

Ma, L. F.; Liu, W. J.; Sheng, Y. P.; Huang, Q. G.; Yang, W. T. J. Appl. Polym. Sci. 2011, 120, 1652–1658. Asandei, A. D.; Simpson, C. P.; Adebolu, O.; Chen, Y. Polym. Prepr. 2011, 52, 554–555. Asandei, A. D.; Moran, I. W.; Saha, G.; Chen, Y. Polym. Mater. Sci. Eng. 2006, 94, 597–598. Asandei, A. D.; Chen, Y. Macromolecules 2006, 39, 7549–7554. Asandei, A. D.; Simpson, C. P. Polym. Prepr. 2008, 49, 75–76. Reardon, D.; Conan, F.; Gambarotta, S.; Yap, G.; Wang, Q. J. Am. Chem. Soc. 1999, 121, 9318–9325. Milione, S.; Cavallo, G.; Tedesco, C.; Grassi, A. J. Chem. Soc. Dalton Trans. 2002, 1839–1846. Colamarco, E.; Milione, S.; Cuomo, C.; Grassi, A. Macromol. Rapid Commun. 2004, 25, 450–454. Lang, J. R. V.; Denner, C. E.; Alt, H. G. J. Mol. Catal. A 2010, 322, 45–49. Shaver, M. P.; Hanhan, M. E.; Jones, M. R. Chem. Commun. 2010, 46, 2127–2129. Allan, L. E. N.; Cross, E. D.; Francis-Pranger, T. W.; Hanhan, M. E.; Jones, M. R.; Pearson, J. K.; Perry, M. R.; Storr, T.; Shaver, M. P. Macromolecules 2011, 44, 4072–4081. Perry, M. R.; Allan, L. E. N.; Decken, A.; Shaver, M. P. Dalton Trans. 2013, 42, 9157–9165. Lee, M.; Minoura, Y. J. Chem. Soc. Faraday Trans. 1978, 74, 1726–1737. Lee, M.; Morigami, T.; Minoura, Y. J. Chem. Soc. Faraday Trans. 1978, 74, 1738–1749. Lee, M.; Utsumi, K.; Minoura, Y. J. Chem. Soc. Faraday Trans. 1979, 75, 1821–1829. Lee, M.; Ishida, Y.; Minoura, Y. J. Polym. Sci. A Polym. Chem. 1982, 20, 457–465. Grishin, D.; Semyonycheva, L.; Artemov, A.; Telegina, E.; Valetova, N.; Illichev, I. Appl. Organomet. Chem. 2003, 17, 717–722. Grishin, D. F.; Valetova, N. B.; Il’ichev, I. S.; Semenycheva, L. L.; Artemov, A. N.; Sazonova, E. V. Polym. Sci., Ser. B 2005, 47, 163–166. Valetova, N. B.; Semyonycheva, L. L.; Illichev, I. S.; Artemov, A. N.; Grishin, D. F. Appl. Organomet. Chem. 2005, 19, 971–974. Valetova, N. B.; Semenycheva, L. L.; Il’ichev, I. S.; Artemov, A. N.; Grishin, D. F. Russ. J. Appl. Chem. 2007, 80, 818–821. Champouret, Y.; Baisch, U.; Poli, R.; Tang, L.; Conway, J. L.; Smith, K. M. Angew. Chem. Int. Ed. 2008, 47, 6069–6072. Champouret, Y.; MacLeod, K. C.; Baisch, U.; Patrick, B. O.; Smith, K. M.; Poli, R. Organometallics 2010, 29, 167–176. Le Grognec, E.; Claverie, J.; Poli, R. J. Am. Chem. Soc. 2001, 123, 9513–9524. Stoffelbach, F.; Poli, R.; Richard, P. J. Organomet. Chem. 2002, 663, 269–276. Stoffelbach, F.; Poli, R.; Maria, S.; Richard, P. J. Organomet. Chem. 2007, 692, 3133–3143. Stoffelbach, F.; Claverie, J.; Poli, R. C. R. Acad. Sci. Paris C 2002, 5, 37–42. Aliwi, S. M.; Bamford, C. H.; Mullik, S. U. J. Polym. Sci. Polym. Symp. 1975, 50, 33–50. Asandei, A. D.; Adebolu, O. I.; Simpson, C. P. J. Am. Chem. Soc. 2012, 134, 6080–6083. Simpson, C. P.; Adebolu, O. I.; Kim, J.-S.; Vasu, V.; Asandei, A. D. Macromolecules 2015, 48, 6404–6420. Morales-Cerrada, R.; Ladmiral, V.; Gayet, F.; Fliedel, C.; Poli, R.; Améduri, B. Polymers 2020, 12, 384/1–384/17. Grishin, I. D.; Krivykh, V. V.; Shchepalov, A. A.; Taits, E. S.; Ustynyuk, N. A.; Grishin, D. F. Russ. Chem. Bull. 2009, 58, 1866–1871. Kanagasabapathy, S.; Serero, D.; Silie, D.; Prost, S.; Ruiz-Guerrero, R.; Claverie, J. Res. Discl. 1998, P1595–P1604. Xue, Z.; Poli, R. J. Polym. Sci. A Polym. Chem. 2013, 51, 3494–3504. Wang, J. R.; Zhou, J.; Sharif, H.; He, D.; Ye, Y. S.; Xue, Z. G.; Xie, X. L. RSC Adv. 2015, 5, 96345–96352. Shaver, M. P.; Allan, L. E. N.; Gibson, V. C. Organometallics 2007, 26, 4725–4730. Poli, R.; Shaver, M. P. Chem. Eur. J. 2014, 20, 17530–17540. Allan, L. E. N.; MacDonald, J. P.; Reckling, A. M.; Kozak, C. M.; Shaver, M. P. Macromol. Rapid Commun. 2012, 33, 414–418. Allan, L. E. N.; MacDonald, J. P.; Nichol, G. S.; Shaver, M. P. Macromolecules 2014, 47, 1249–1257. Schroeder, H.; Lake, B. R. M.; Demeshko, S.; Shaver, M. P.; Buback, M. Macromolecules 2015, 48, 4329–4338. Coward, D. L.; Lake, B. R. M.; Shaver, M. P. Organometallics 2017, 36, 3322–3328. Kato, M.; Kamigaito, M.; Sawamoto, M.; Higashimura, T. Macromolecules 1995, 28, 1721–1723. Takahashi, H.; Ando, T.; Kamigaito, M.; Sawamoto, M. Macromolecules 1999, 32, 3820–3823. Braunecker, W. A.; Itami, Y.; Matyjaszewski, K. Macromolecules 2005, 38, 9402–9404. Braunecker, W. A.; Brown, W. C.; Morelli, B.; Tang, W.; Poli, R.; Matyjaszewski, K. Macromolecules 2007, 40, 8576–8585. Wayland, B. B.; Poszmik, G.; Mukerjee, S. J. Am. Chem. Soc. 1994, 116, 7943–7944. Arvanatitopoulos, L. D.; Greuel, M. P.; Harwood, H. J. Polym. Prepr. 1994, 35, 549–550. Arvanitopoulos, L. D.; Greuel, M. P.; King, B. M.; Shim, A. K.; Harwood, H. J. ACS Symp. Ser. 1998, 685, 316–331. Lu, Z.; Fryd, M.; Wayland, B. B. Macromolecules 2004, 37, 2686–2687. Wayland, B. B.; Basickes, L.; Mukerjee, S.; Wei, M.; Fryd, M. Macromolecules 1997, 30, 8109–8112. Wayland, B. B.; Mukerjee, S.; Poszmik, G.; Woska, D. C.; Basickes, L.; Gridnev, A. A.; Fryd, M.; Ittel, S. D. ACS Symp. Ser. 1998, 685, 305–315. Wayland, B. B.; Peng, C.-H.; Fu, X.; Lu, Z.; Fryd, M. Macromolecules 2006, 39, 8219–8222. Wayland, B. B.; Fu, X.; Peng, C.-H.; Lu, Z.; Fryd, M. ACS Symp. Ser. 2006, 944, 358–371. Peng, C.-H.; Fryd, M.; Wayland, B. B. Macromolecules 2007, 40, 6814–6819. Zhao, Y.; Dong, H.; Li, Y.; Fu, X. Chem. Commun. 2012, 48, 3506–3508. Dong, H. L.; Hou, T. J.; Zhao, Y. G.; Fu, X. F.; Li, Y. Y. Comput. Theor. Chem. 2012, 1001, 51–59. Peng, C. H.; Scricco, J.; Li, S.; Fryd, M.; Wayland, B. B. Macromolecules 2008, 41, 2368–2373. Li, S.; de Bruin, B.; Peng, C. H.; Fryd, M.; Wayland, B. B. J. Am. Chem. Soc. 2008, 130, 13373–13381. Hsu, C. S.; Yang, T. Y.; Peng, C. H. Polym. Chem. 2014, 5, 3867–3875. Maria, S.; Kaneyoshi, H.; Matyjaszewski, K.; Poli, R. Chem. Eur. J. 2007, 13, 2480–2492. Zhao, Y. G.; Yu, M. M.; Fu, X. F. Chem. Commun. 2013, 49, 5186–5188. Zhao, Y.; Yu, M.; Zhang, S.; Liu, Y.; Fu, X. Macromolecules 2014, 47, 6238–6245. Debuigne, A.; Caille, J. R.; Detrembleur, C.; Jerome, R. Angew. Chem. Int. Ed. 2005, 44, 3439–3442. Debuigne, A.; Caille, J. R.; Jérôme, R. Angew. Chem. Int. Ed. 2005, 44, 1101–1104. Bryaskova, R.; Detrembleur, C.; Debuigne, A.; Jerome, R. Macromolecules 2006, 39, 8263–8268. Santhosh Kumar, K. S.; Gnanou, Y.; Champouret, Y.; Daran, J.-C.; Poli, R. Chem. Eur. J. 2009, 15, 4874–4885. Kaneyoshi, H.; Matyjaszewski, K. Macromolecules 2005, 38, 8163–8169. Banerjee, S.; Bellan, E. V.; Gayet, F.; Debuigne, A.; Detrembleur, C.; Poli, R.; Améduri, B.; Ladmiral, V. Polymers 2017, 9, 702. Bellan, E. V.; Thevenin, L.; Gayet, F.; Fliedel, C.; Poli, R. ACS Macro Lett. 2017, 6, 959–962. Kaneyoshi, H.; Matyjaszewski, K. Macromolecules 2006, 39, 2757–2763. Debuigne, A.; Willet, N.; Jerome, R.; Detrembleur, C. Macromolecules 2007, 40, 7111–7118. Debuigne, A.; Schoumacher, M.; Willet, N.; Riva, R.; Zhu, X. M.; Rutten, S.; Jerome, C.; Detrembleur, C. Chem. Commun. 2011, 47, 12703–12705. Debuigne, A.; Morin, A. N.; Kermagoret, A.; Piette, Y.; Detrembleur, C.; Jérôme, C.; Poli, R. Chem. Eur. J. 2012, 18, 12834–12844. Hurtgen, M.; Liu, J.; Debuigne, A.; Jerome, C.; Detrembleur, C. J. Polym. Sci. A Polym. Chem. 2012, 50, 400–408.

Reversible Homolysis of Metal-Carbon Bonds

462. 463. 464. 465. 466. 467. 468. 469. 470. 471. 472. 473. 474. 475. 476. 477. 478. 479. 480. 481. 482. 483. 484. 485. 486. 487. 488. 489. 490. 491. 492. 493. 494. 495. 496. 497. 498. 499. 500. 501. 502. 503. 504.

85

Piette, Y.; Debuigne, A.; Jérôme, C.; Bodart, V.; Poli, R.; Detrembleur, C. Polym. Chem. 2012, 3, 2880–2891. Banerjee, S.; Ladmiral, V.; Debuigne, A.; Detrembleur, C.; Poli, R.; Améduri, B. Angew. Chem. Int. Ed. 2018, 57, 2934–2937. Falireas, P. G.; Ladmiral, V.; Debuigne, A.; Detrembleur, C.; Poli, R.; Ameduri, B. Macromolecules 2019, 52, 1266–1276. Kermagoret, A.; Debuigne, A.; Jerome, C.; Detrembleur, C. Nat. Chem. 2014, 6, 179–187. Bryaskova, R.; Willet, N.; Degee, P.; Dubois, P.; Jerome, R.; Detrembleur, C. J. Polym. Sci. A Polym. Chem. 2007, 45, 2532–2542. Kermagoret, A.; Wenn, B.; Debuigne, A.; Jerome, C.; Junkers, T.; Detrembleur, C. Polym. Chem. 2015, 6, 3847–3857. Demarteau, J.; Améduri, B.; Ladmiral, V.; Mees, M. A.; Hoogenboom, R.; Debuigne, A.; Detrembleur, C. Macromolecules 2017, 50, 3750–3760. Debuigne, A.; Warnant, J.; Jerome, R.; Voets, I.; de Keizer, A.; Stuart, M. A.; Detrembleur, C. Macromolecules 2008, 41, 2353–2360. Debuigne, A.; Michaux, C.; Jérôme, C.; Jérôme, R.; Poli, R.; Detrembleur, C. Chem. Eur. J. 2008, 14, 7623–7637. Detrembleur, C.; Versace, D. L.; Piette, Y.; Hurtgen, M.; Jerome, C.; Lalevee, J.; Debuigne, A. Polym. Chem. 2012, 3, 1856–1866. Debuigne, A.; Jerome, C.; Detrembleur, C. Angew. Chem. Int. Ed. 2009, 48, 1422–1424. Debuigne, A.; Poli, R.; De Winter, J.; Laurent, P.; Gerbaux, P.; Dubois, P.; Wathelet, J.-P.; Jérôme, C.; Detrembleur, C. Chem. Eur. J. 2010, 16, 1799–1811. Debuigne, A.; Poli, R.; De Winter, J.; Laurent, P.; Gerbaux, P.; Wathelet, J.-P.; Jérôme, C.; Detrembleur, C. Macromolecules 2010, 43, 2801–2813. Debuigne, A.; Champouret, Y.; Jérôme, R.; Poli, R.; Detrembleur, C. Chem. Eur. J. 2008, 14, 4046–4059. Debuigne, A.; Poli, R.; Jérôme, R.; Jérôme, C.; Detrembleur, C. ACS Symp. Ser. 2009, 1024, 131–148. Wang, F. S.; Yang, T. Y.; Hsu, C. C.; Chen, Y. J.; Li, M. H.; Hsu, Y. J.; Chuang, M. C.; Peng, C. H. Macromol. Chem. Phys. 2016, 217, 422–432. Wang, Z.; Poli, R.; Detrembleur, C.; Debuigne, A. Macromolecules 2019, 52, 8976–8988. Gridnev, A. A. Polym. J. 1992, 24, 613–623. Heuts, J. P. A.; Forster, D. J.; Davis, T. P. Macromol. Rapid Commun. 1999, 20, 299–302. Roberts, G. E.; Davis, T. P.; Heuts, J. P. A.; Russell, G. T. J. Polym. Sci. A Polym. Chem. 2002, 40, 782–792. Engelis, N. G.; Anastasaki, A.; Whitfield, R.; Jones, G. R.; Liarou, E.; Nikolaou, V.; Nurumbetov, G.; Haddleton, D. M. Macromolecules 2018, 51, 336–342. Ng, Y. H.; di Lena, F.; Chai, C. L. L. Macromol. Res. 2012, 20, 473–476. Sherwood, R. K.; Kent, C. L.; Patrick, B. O.; McNeil, W. S. Chem. Commun. 2010, 46, 2456–2458. Liao, C.-M.; Hsu, C.-C.; Wang, F.-S.; Wayland, B. B.; Peng, C.-H. Polym. Chem. 2013, 4, 3098–3104. Zhao, Y.; Yu, M.; Zhang, S.; Wu, Z.; Liu, Y.; Peng, C.-H.; Fu, X. Chem. Sci. 2015, 6, 2979–2988. Zhao, Y. G.; Zhang, S. L.; Wu, Z. Q.; Liu, X.; Zhao, X. Y.; Peng, C. H.; Fu, X. F. Macromolecules 2015, 48, 5132–5139. Wu, Z. Q.; Wang, Z. K.; Wang, B. W.; Peng, C. H.; Fu, X. F. Macromolecules 2020, 53, 212–222. Wang, Y.; Zhao, Y.; Zhu, S.; Zhou, X.; Xu, J.; Xie, X.; Poli, R. Angew. Chem. Int. Ed. 2020, 59, 5988–5994. Chiang, L.; Allan, L. E. N.; Alcantara, J.; Wang, M. C. P.; Storr, T.; Shaver, M. P. Dalton Trans. 2014, 43, 4295–4304. Langlotz, B. K.; Loret Fillol, J.; Gross, J. H.; Wadepohl, H.; Gade, L. H. Chem. Eur. J. 2008, 14, 10267–10279. Bagchi, V.; Raptopoulos, G.; Das, P.; Christodoulou, S.; Wang, Q. W.; Ai, L.; Choudhury, A.; Pitsikalis, M.; Paraskevopoulou, P.; Stavropoulos, P. Polyhedron 2013, 52, 78–90. Santhosh Kumar, K. S.; Li, Y.; Gnanou, Y.; Baisch, U.; Champouret, Y.; Poli, R.; Robson, K. C. D.; McNeil, W. S. Chem. Asian J. 2009, 4, 1257–1265. Chen, Y. H.; Chen, S. J.; Li, J. Q.; Wu, Z. Q.; Lee, G. H.; Liu, Y. H.; Cheng, W. T.; Yeh, C. Y.; Peng, C. H. J. Polym. Sci. 2020, 58, 101–113. Wayland, B. B.; Poszmik, G.; Fryd, M. Organometallics 1992, 11, 3534–3542. Ribelli, T. G.; Fantin, M.; Daran, J.-C.; Augustine, K. F.; Poli, R.; Matyjaszewski, K. J. Am. Chem. Soc. 2018, 140, 1525–1534. Ribelli, T. G.; Matyjaszewski, K.; Poli, R. J. Coord. Chem. 2018, 71, 1641–1668. Matyjaszewski, K.; Woodworth, B. E. Macromolecules 1998, 31, 4718–4723. Schröder, K.; Konkolewicz, D.; Poli, R.; Matyjaszewski, K. Organometallics 2012, 31, 7994–7999. Ribelli, T. G.; Rahaman, S. M. W.; Daran, J.-C.; Krys, P.; Matyjaszewski, K.; Poli, R. Macromolecules 2016, 49, 7749–7757. Wang, Y.; Soerensen, N.; Zhong, M.; Schroeder, H.; Buback, M.; Matyjaszewski, K. Macromolecules 2013, 46, 683–691. Rahaman, S. M. W.; Matyjaszewski, K.; Poli, R. Polym. Chem. 2016, 7, 1079–1087. Ribelli, T. G.; Augustine, K. F.; Fantin, M.; Krys, P.; Poli, R.; Matyjaszewski, K. Macromolecules 2017, 50, 7920–7929. Ribelli, T. G.; Rahaman, S. M. W.; Matyjaszewski, K.; Poli, R. In Reversible Deactivation Radical Polymerization: Mechanisms and Synthetic Methodologies; Tsarevsky, N., Gao, H., Matyjaszewski, K., Sumerlin, B., Eds.; American Chemical Society: Washington DC, 2018; vol. 1284; pp 135–159. 505. Ribelli, T. G.; Lorandi, F.; Fantin, M.; Matyjaszewski, K. Macromol. Rapid Commun. 2019, 40, 1800616. 506. Enciso, A. E.; Lorandi, F.; Mehmood, A.; Fantin, M.; Szczepaniak, G.; Janesko, B. G.; Matyjaszewski, K. Angew. Chem. Int. Ed. 2020, 59, 14910–14920.

1.04

Very Low Oxidation States in Organometallic Chemistry

C Gunnar Werncke, Chemistry Department, Philipps-University Marburg, Marburg, Germany © 2022 Elsevier Ltd. All rights reserved.

1.04.1 Preface 1.04.1.1 Main-group (organo)metal polyanions 1.04.1.2 Mononuclear metal anions 1.04.2 Organometallic compounds with negative oxidation states of the transition metal 1.04.2.1 Comment on oxidation state formalism and redox non-innocence 1.04.2.2 Carbonyl metallates 1.04.2.2.1 Reactivity of anionic carbonyl metallates 1.04.2.3 Isonitrile-based metallates 1.04.2.4 Alkene metallate complexes 1.04.2.4.1 Homoleptic arene-based metallate complexes 1.04.2.4.2 Synthesis and structure of arene metallates 1.04.2.4.3 Reactivity 1.04.2.5 Carbene-based metallates 1.04.2.6 Honorable mentions 1.04.3 Concluding remarks and outlook Acknowledgement References

86 87 88 88 88 89 91 93 95 96 97 99 101 102 103 104 104

Nomenclature 18c6 AM anth cod COT Cp crypt-2.2.2 diphen en HMPA i Pr Mes naph PPN QTAIM thf tmeda VE xyl

1.04.1

18-Crown-6 (1,4,7,10,13,16-hexaoxacyclooctadecane) Alkali metal Anthracene 1,5-Cyclooctadiene 1,3,5,7-Cyclooctatetraene Cyclopentadienyl 4,7,13,16,21,24-Hexaoxa-1,10-diazabicyclo[8.8.8]hexacosane Diphenyl Ethylenediamine/ethane-1,2-diamine Hexamethylphosphoramide iso-Propyl 1,3,5-Trimethylphenyl Naphthalene Bis(triphenylphosphine)iminium Quantum theory of atoms in molecules Tetrahydrofuran N,N,N0 ,N0 -Tetramethylethane-1,2-diamine Valence electrons 2,6-Dimethylphenyl

Preface

In classic coordination and related organometallic chemistry, a metal complex consists of a metal cation surrounded by mainly X-type (anionic) as well as additional L-type (neutral) ligands. For each metal a certain (range of ) oxidation state(s) is commonly found, e.g. +IV for the 4th group (Ti, Zr, Nb), +II/+III for the 8th group (Fe, Ru) or +I for the 11th group. With these differences in mind, the term “low-valent” usually refers to oxidation states below these respective numbers, and its meaning naturally differs by group and metal. For groups 4 and 5 (and to a lesser extent group 3) the term low-valent is commonly attributed to the oxidation states 0 to +II and 0/+I for group 6–10. For group 11 (Cu, Ag and Au) this terminology is less used, and refers solely to oxidation state 0 because the +1 oxidation state is so common. Zerovalent group 11 complexes are scarce due to their tendency for metal-metal interactions, aggregation and metal cluster formation.1 For group 12 (Zn, Cd, Hg), +I is deemed as low-valent, especially for zinc.

86

Comprehensive Organometallic Chemistry IV

https://doi.org/10.1016/B978-0-12-820206-7.00007-X

Very Low Oxidation States in Organometallic Chemistry

87

Building on this definition of “low-valent” in terms of oxidation state, very low oxidation states thus refers to complexes where the metal is found in even lower, formally negative oxidation state. Compounds with metal in a negative oxidation state or even “naked” metal (poly) anions have been known for more than a century.2,3 Most studies address carbonyl-supported metallates of transition metals like Na2[Fe(CO)4] or main group polyanions (namely Zintl anions such as [Pb9]4−). Apart from these, very low-valent organometallic metallates are a rather underused class of molecules but offer opportunities for small molecule activation and for fundamental study. As such the aim of the chapter is to give a broad overview of the properties, reactivity and of valent metal compounds with a focus on transition metal complexes. This is intended to inspire organometallic and coordination chemists to take up the remarkable work in the area of transition metal compounds in very low oxidation states that was pioneered by scientific giants such as Hieber, Collman, Ellis and others.

1.04.1.1

Main-group (organo)metal polyanions

The chemistry of metal anions and organometallic compounds with the metal ion in a negative oxidation state was initiated more than 100 years ago by M. Joannis. He reported in 1891 on deeply green and wine-red coloured solutions upon reaction of sodium metal with lead, mercury and antimony, respectively. In the following decades, work by Kraus, Peck, Smyth and Zintl expanded this to other main group metals and led to the discovery of highly charged homometallic polyanions such as [Pb9]4− or [Sb7]3− (Scheme 1).4–9 These polyanions could also be obtained upon dissolution of intermetallic phases such as Na4Pb9 in liquid ammonia, connecting salt-like phases with soluble polyanions.

Scheme 1 Synthesis of a Zintl phase (Na4Pb9) and extraction of a molecular Zintl polyanion.

These phases can generally be obtained via melting stoichiometric amounts of the alkali and main group metals. The chemistry of these so-called Zintl phases and derived Zintl anions has since evolved into an ever-growing research field. The structural characterisation of discrete metallic polyanions was accelerated by the switch of the solvent from NH3 to more convenient solvents such as ethylenediamine or dimethylformamide. A strong donor capability of the solvent is necessary to stabilize the polyanions in solution as well as to separate it from the alkali counterions. However, the composition of the polyanion in the Zintl phases is not necessarily maintained in solution upon extraction. A further, important synthetic step was the use of crypt-2.2.2 (4,7,13,16,21,24hexaoxa-1,10-diazabicyclo[8.8.8]hexacosane) as an alkali metal chelating agent, which was popularized by Corbett in the 1970s and facilitated the crystallographic characterisation and isolation of salts containing these polyanions. To a lesser degree, this was also expanded to the cheaper 18-crown-6 (1,4,7,10,13,16-hexaoxacyclooctadecane) and cation exchange with NMe+4. The identity of the specific counterion can thereby influence the composition of the metal polyanions. The synthesis of Zintl anions could be extended to binary polyanions (e.g. [Sb2Sn2]2− or [TlSn9]3−). These endeavours led to the development of the Zintl-KlemmBussman or pseudo-element concept (Fig. 1 left), which allows conceptual connection of the structure of a polyanion to the one made out of the element that possesses the same count of valence electrons (e.g. tetrahedral P4 and tetrahedral [GaBi3]2− or [Sb2Bi2]2−). Intriguingly, Zintl anions can further incorporate transition metal ions forming intermetalloid (endohedral) cluster anions with a metal encapsulated within the polyanion (e.g. [Ni@Ge9]3−, Fig. 1 right).

Fig. 1 Structural interconnection between P4 and a binary Zintl anion (left), as well as formation of an intermetalloid cluster (right) (Pn ¼ pnictogene/group 15 element; Tr ¼ triel/group 13 element).

These examples are all “naked” polyanions with metal ions (partially) in a subvalent state without any organometallic ligation. However, they can be starting points for organometallic cluster derivatives bearing one or several metal-carbon bonds. Three principal synthetic pathways are employed: (A) direct functionalisation using an organohalide or an organometallic main-group halide R–X in a classic salt-metathesis reaction giving a capping “MR” unit; (B) the reaction with (silylated) alkynes leading to vinylation of the cluster7; (C) introduction of M(CO)3 fragments (M ¼ Cr, Mo, and W), Ir(cod)+, CuPR+3 (R ¼ iPr, Cy) or MR+ (M ¼ Zn, Cd and R ¼ C6H5, iPr, Mes and Sn(alkyl)3) units. The main group metal polyanions thereby act as an inorganic ligand. For comprehensive insights into polyatomic Zintl anions and their chemistry the reader is referred to inspiring accounts from

88

Very Low Oxidation States in Organometallic Chemistry

Corbett,10,11 Sevov,12 Fässler,13 and Dehnen.14,15 These cover organofunctionalized clusters derived from Zintl anions which technically classify as subvalent metal containing organometallic compounds, but which will not be discussed further here.

1.04.1.2

Mononuclear metal anions

Monoatomic metal anions were detected first in the gas phase by mass spectrometry, with the first reports on Hg− in the 1930s during endeavours to experimentally determine electron affinities of elements and small molecules.16–19 The actual formation of Hg− in these reports remains in doubt, as it is calculated as unfeasible.20 In any case, these first studies already showed that some alkali and transition metals can have surprisingly high electron affinities (EA) which was supported by computational methods. For example, the highest EAs of all metals are possessed by Au (2.31 eV) and Pt (2.13 eV), which even surpass those of chalcogenides (S: 2.08).21 The first unambiguously observed metal anion in the gas phase was Li− in 1947,22 and was later expanded to nearly all main group and transition metal elements.23 The Lewis basic character of alkali metal monoanions was demonstrated through adduct formation with BH3 in the gas phase (described as Na−!BH3), but the bonding in this adduct is still the subject of controversy.24–27 In condensed phase, monometallic anions M− have been observed in solid state or in solution for Pt, Ag, Au and the heavier alkali metals (Na–Cs).28,29 The first proposed isolation of a monoatomic metal anion was for CsAu in the 1950’s in solid state, which was followed by the detection of Au− in NH3 solutions. The first example of a structurally verified auride complex was [Me4N][Au] (Scheme 2). The auride anion is thought to behave as a pseudohalide although its coordination chemistry remains unexplored (vide infra).30

Scheme 2 Isolable complexes/compounds with unligated monometalic anions (am ¼ solvated in NH3).

Another remarkable example is the disproportionation of sodium metal by addition of crypt-2.2.2 to give Na{crypt-2.2.2}[Na] (Scheme 2).31 This could also be expanded to salts of the heavier alkali metals.32 The recent syntheses of solid Cs2Pt and Ba2Pt from the elements are a further milestone (Scheme 2). Thereby Pt2− constitutes the first and only example of an isolable “ligand free” metal dianion which is a formal transition metal analogue of a chalcogenide dianion.33,34

1.04.2

Organometallic compounds with negative oxidation states of the transition metal

1.04.2.1

Comment on oxidation state formalism and redox non-innocence

The chemistry of monometallic organometallic compounds with the transition metal in a negative oxidation state is so far restricted to the d-block elements. Before going deeper into the respective chemistry, a note on the attribution of an oxidation state to the metal is needed. Usually the determination of the metal’s formal oxidation state is achieved by conceptually removing all ligands in their most stable (closed shell) forms. Examples of such ligands in organometallic chemistry are halides, hydrides, carbon monoxide, phosphines, phosphites, (iso)nitriles, and carbenes. With p-alkene and p-alkyne complexes, the complexes have an alternative resonance structure as a metallacyclopropane or metallacyclopropene respectively, where the metal formally adopts an oxidation state that is two higher. In terms of reactivity metallacyclopropane/ene complexes can release the alkene/alkyne and can thus be viewed as masked lower valent compounds.35–39 The formal approach of assigning metal-ligand bonding electrons to the ligand neglects the fact that the electronic structure of free and coordinated ligands is greatly influenced by the interaction with the metal. This holds especially true for ligands that can easily accept or release electrons, and for which the complexes may be reasonably described as metal stabilized radicals or radical anions, with the metal oxidation state changing accordingly. To address these complications C. K. Jorgensen established the terms innocence and non-innocence.40 Ligands that are innocent are those that enable an accurate assignment of oxidation states within a complex, whereas non-innocent ligands are more ambiguous. Prime examples of non-innocent ligands are in reduced bipyridine complexes such as [M(bipy)3]n (Sc, Y, Ti, Zr, Hf, V, Nb, Ta; n ¼ 1+, 0, 1−, 2−, 3−), in which a formal low or negative oxidation state was initially assigned to the metal, based on the usual oxidation state formalisms. However, it has been shown that these complexes are more accurately described as having the metal ions in a higher oxidation state (+III for Sc, Y, Ti; +IV for Zr, Hf, V, Nb; +V for Ta) with the redox chemistry taking place primarily on the ligand.41 The situation is similar for complexes with the analogous biphosphinine ligands.42 Another example is the nitrosyl ligand, which can be regarded as anionic (NO−), neutral (NO) or cationic (NO+) depending on the nature of the interaction with the coordinating metal. Though it is often possible to distinguish NO− complexes from NO+ complexes based on the geometry (bent vs. linear), the assignment of the oxidation states of a metal bound NO ligand can be ambiguous, especially for polynitrosyl compounds.

Very Low Oxidation States in Organometallic Chemistry

89

Therefore, Enemark and Feltham introduced the notation of {M(NO)n}m (n ¼ number of nitrosyl ligands; m ¼ number of unpaired electrons with p-symmetry) which does not localize oxidation states to the metal and the ligand but rather gives the sum of d electrons on the metal and electrons in the p-system of the nitrosyl: thus it treats the metal nitrosyl fragment as a whole.43 In the last two decades, chemists such as Kaim and Wieghardt documented examples in which a variety of ligands can act in an non-innocent fashion. For example low-lying p orbitals can be (partially) filled under highly reducing conditions, e.g. (di)imines, (di)iminopyridines and more recently N-heterocyclic carbenes or cyclic alkyl amino carbenes.44–50 Thereby bond metrics proved to be usefel to deduce the oxidation state of certain ligands (e.g. for 2,20 -bipyridine or iminopyridines).41,51 Further more sophisticated analytical (e.g. X-ray absorption spectroscopy) and computational techniques (CASSCF) become more available that allow for deeper insights into bonding situations.52–57 This is especially needed for systems that carry pronounced multireference character,24,58–60 such as paramagnetic compounds that are quasi-degenerate due to low-lying excited states. Further, the description of the metal-ligand interaction in terms of dative and electron sharing bonds can thereby change when going from regular Werner-type compounds to more unusual molecules (e.g. the observation of an inverted ligand field for a formal copper(III) imide which is better described as copper(I) singlet nitrene complex).57,61–65 As another example, alkyne complexes with a low-valent metal ion should not only be regarded as p-complexes or metallacyclopropenes as established via the Chatt-DewarDuncanson-model but can share some 3c-3e (3 centre-3 electron) character with radical character of the alkyne.66 The description of compounds bearing the metal in negative oxidation states has to be treated with even more care, as many reports date back decades when structural and electronic insights often relied solely on IR spectroscopy and lacked structural metrics, and when the consideration of redox non-innocence was less common. Irrespective of the accuracy of the oxidation state formalism in these complexes, in the context of this chapter it is still useful to understand their stability and reactivity. A renewed interest into subvalent compounds with contemporary analytical and computational methods may reveal further non-innocent ligand character in so far “unsuspicious” ligands in such complexes in very low oxidation states. Overall, the consideration of ligand redox activity is inevitable in any description of complexes having a (formally) negatively charged metal ion. The isolation of such substances is dependent on supporting ligands that are capable of metal-to-ligand backbonding, and thus partially alleviate the electron density of the metal. At the same time, these supporting ligands must be tolerant of the strongly reducing conditions. In this regard carbon monoxide, isonitriles, alkenes, arenes and to a lesser extent PF3, carbenes and phosphites have found heavy use. In the following, we will concentrate on the organometallic derivatives.

1.04.2.2

Carbonyl metallates

Historically, the chemistry of organometallics in negative oxidation states started with the use of compounds bearing carbonyl ligands (from now on coined as carbonyl metallates). CO is a strong field ligand and is known to stabilize metal ions in low oxidation states by synergistic carbonyl-metal s-donation and more importantly metal-to-ligand p-backbonding.67–69 The availability of zerovalent, homoleptic metal carbonyls provided a starting point for the organometallic chemistry with central metal anions. The earliest formulation of anionic carbonyl metallates dates back to 1929 when Hock and Stuhlmann speculated on the formation of Hg[Fe(CO)4] upon reaction of Fe(CO)5 with mercury salts.70 In 1931, Hieber described the reaction of Fe(CO)5 and Co2(CO)9 with bases such as Ba(OH)2 or NaOH (Hieber thus named it the “Basen-Reaktion” or “Base reaction”), which led to the identification of the first volatile transition metal hydrides.71 In this report Hieber already commented on the possibility of the formation of an [Fe(CO)4]2− anion (Scheme 3), but he did not found evidence at that time. This hypothesis was finally substantiated by Krumholz et al. via isolation of the dipotassium salt as colourless crystals from an aqueous, deoxygenated KOH solution.72 The reductant in these reactions is CO, which is oxidized to carbonate. Remarkably, Dewar already described the formation of colourless crystals from an identical reaction nearly 30 years prior in 1905,73 which might be actually the first observation of K2[Fe(CO)4] although he was unaware of its constitution at that time.

Scheme 3 Hieber’s Base reaction for the synthesis of the carbonyl ferrate [Fe(CO)4]2−.

The formation of carbonyl metallates could also be extended to dinuclear ([Fe2(CO)8]2−) or even trinuclear dianions ([Fe2(CO)11]2−) by using Fe2(CO)9 and Fe3(CO)12, respectively.74,75 The corresponding reactions of the carbonyl compounds of the neighbouring elements such as [Mn2(CO)10] and [Co2(CO)10] give the monomeric metallates [Mn(CO)5]− and [Co(CO)4]−.76–78 As the M:CO ratio remains unchanged in the neutral and anionic carbonyl complexes for these odd-numbered elements, the reaction is preceded by a redox disproportionation of the metal carbonyl into the carbonyl metallate(−I), M2+ and free CO.79 For metal carbonyls like Cr(CO)6 the “Basen-Reaktion” is unsuccessful, due to the highly reactive nature of the formed [Cr(CO)5]2− dianion.80

90

Very Low Oxidation States in Organometallic Chemistry

When Hieber reacted organic bases (e.g. pyridine or ethylenediamine (en)) with metal(0) carbonyls under anhydrous conditions, redox disproportionation was regularly observed. For example, gradual heating of Fe3(CO)9 with ethylenediamine yielded first [Fe(en)3][Fe2(CO)11], then [Fe(en)3][Fe2(CO)8] and at 145  C ultimately [Fe(en)3][Fe(CO)4]. Although the use of Lewis bases gave a variety of carbonyl metallates, e.g. [V(CO)6]−, [Ni2(CO)8]2−, or [Ni4(CO)9]2−, the parallel formation of a metal containing cation is not practical in terms of yield and product isolation.79,81 The most frequently used synthesis of carbonyl metallates in current times is the reduction of the homoleptic metal carbonyls with an alkali metal AM (Scheme 4). Historically, alkali metal amalgams or finely dispersed alkali metals were used, with mixed results. This was extended to liquid ammonia or hexamethylphosphoramide (HMPA), in which alkali metals can be dissolved. Later, the use of alkali naphthalenides was introduced as a softer reducing agent.82–85

Scheme 4 Synthetic approaches to carbonyl based metallates.

In certain cases, the direct reduction of homoleptic metal halides failed. For that the use of Lewis base adducts of metal carbonyls (e.g. (M(CO)4dmpe] or higher valent carbonyl metal halides (e.g. [M(CO)m(Cl)n]) proved to be successful. This introduces labile ligands that are more likely to be extruded upon reduction. The success of each synthetic approach is dependent on the metal carbonyl, the envisioned negative oxidation state and further substituents. The synthesis of reduced carbonyl metallates can be accompanied by alkali metal mediated reduction of CO, stemming from using a CO atmosphere or liberation upon reduction, which gives the explosive di(alkali metal) acetylenediolate (AM)2C2O2. Further, some metallates themselves are shown to be not only highly reactive but can also decompose in a violent fashion (e.g. [Nb(CO)4]3− or [Ta(CO)4]3−).86 Nowadays of all group 4-9 metal carbonyl mono- (odd-numbered elements) or dianions (even-numbered elements) are synthetically accessible. For group 5–7 and 9 metals (except Tc) this holds true also for the for tri- and tetraanionic complexes respectively, whose synthesis rely on the use of low temperatures and solvents such as HMPA or liquid ammonia. Their molecular structure is so far unknown and elucidation of their identity is based on IR spectroscopy, elemental analysis as well as derivatisation with organotin chlorides (Table 1).

Table 1

Isolable homoleptic mononuclear carbonyl metallates of group 4–9 (L ¼ CO).

4

5

6

7

8

9

10

[TiL6]2−87

[VL6]−88 [VL5]3−89 [NbL6]−97,98 [NbL5]3−98,99 [TaL6]−97 [TaL5]3−99

[CrL5]2−90–92 [CrL4]4−93 [MoL5]2−90,92 [MoL4]4−93 [WL5]2−90,92 [WL4]4−93

[MnL5]1−77 [MnL4]3−94 [TcL5]−100

[FeL4]2−72

[NiL3]2−

[ReL4]−103 [ReL4]3−94

[OsL4]2−104

[CoL4]−78,95 [CoL3]3−96 [RhL4]−102 [RhL3]3−96 [IrL4]−89 [IrL4]3−96,105

[ZrL6]2−87 [HfL6]2−87

[RuL4]2−101

/ /

Homoleptic group 3 carbonyl metallates as well as those of group 10 and beyond were mostly only detected in the gas phase.106–108 For the former their presumed high reducing capabilities likely leads to irreversible reduction of carbonyl ligands, whereas for late transition metals the metal carbonyl bond becomes rather weak due to the inability of the low-energy d-orbitals to engage in backbonding.109,110 Nonetheless, the first crystal structure of a trigonal planar carbonyl metallate was recently reported, namely [Ni(CO)3]2−, by an fortuitous reaction of a Zintl phase with [(PPh3)2Ni(CO)2].111 The initial insights into the electronic structure of these complexes were from Blanchard, who introduced the concept of negative oxidation state for transition metal carbonylates.112 Based on counting the valencies he further predicted that [Fe(CO)4]2− and [Co(Co)4]− should be isostructural to the stable Ni(CO)4 as they all possess the same electron configuration (18 valence electrons, Fig. 2 top). This was later confirmed by X-ray diffraction analysis.113,114

Very Low Oxidation States in Organometallic Chemistry

91

Fig. 2 Interconnection between isostructural metal carbonyl compounds as well as isoelectronic manganese carbonyls.

The approach of evaluating the reactivity and stability of metal carbonyl fragments via the 18-electron rule applies naturally to anionic metallates as well. This was exploited to rationalize the stability of highly anionic compounds such as [Mn(CO)4]3− and even [Cr(CO)4]4−.94 In the isostructural tetracarbonyl complexes [M(CO)4]n− the effect of the negative charge can be probed via IR spectroscopy (Fig. 2). The v~(CO) values decrease by 150–160 cm−1 per charge in a nearly linear fashion for n ¼ 0, 1, 2.87 The redshift can be attributed to enhanced p-backbonding of the more reduced metal species to the carbonyl ligand with similar s-donation from the ligand to metal.115,116 For [Mn(CO)4]3− the decrease of v~(CO) is smaller, which was attributed to a strong interaction of the tetracarbonyl manganate(−III) with the alkali metal ions.117 The cation-metallate interaction is a common feature in case of unmasked alkali metal cations and can be observed in solution as well solid state. A tight coordination of the cation to the carbonyl ligands can further result in a distortion of the metal carbonyl as evidenced by varying carbonyl IR signatures of the same carbonyl metallate depending on the identity of the alkali metal and its speciation in solution.118,119 For highly reduced metallates with unsolvated alkali metal cations, close contacts with carbonyl complexes are suggested by very red-shifted carbonyl v~(CO), e.g. 1460 cm−1 for Na4[Cr(CO)4] that might be better formulated as [Cr(CONa)4]x.93 Within a series of metal carbonyls reduction by 2e− is balanced by loss of a carbonyl ligand to maintain the 18-electron count. This is exemplified for the [Mn(CO)6−n]+1−2n series (n ¼ 0, 1, 2; Fig. 2 bottom).120 Besides these mononuclear compounds, a few polynuclear carbonyl metallates are also known. These are typically obtained by reduction of the respective oligonuclear metal carbonyls and are exclusively found in a dianionic form and for group 6, 8, 10 metals (e.g. [M2(CO)8]2− and [M3(CO)11]2− (Fe, Ru, Os), [W2(CO)10]2−, or [Ni2(CO)6]2−).75,121–124

1.04.2.2.1

Reactivity of anionic carbonyl metallates

The structure and reactivity of the carbonyl metallates can be understood using the isolobal principle.125 For example, the highly reactive Mn(CO)5 fragment (17 VE) is isolobal with the methyl radical, and readily undergoes dimerization or other radical pathways.125 In contrast, the [Mn(CO)5]− anion, isolobal with the methyl anion, is isolable and reacts readily with electrophiles. The nucleophilicity of the mono- and dianionic metallates is their most dominant reactivity characteristic (Scheme 5). It encompasses the reaction with protons, organohalides, main group halides and transition metal halides (especially those of manganese, cobalt and iron).81,126–131

Scheme 5 Selected reactions of [Fe(CO)4]2− with electrophiles.

92

Very Low Oxidation States in Organometallic Chemistry

The most remarkable dianionic carbonyl metallates are those of the group 4 metals (titanium, zirconium and hafnium) which are thermally unstable.87,132 A titanium example gives insights into their reactivity (Scheme 6). Treatment of [Ti(CO)6]2− with Ph3PAuCl gave a compound with a gold-titanium bond whereas its reaction with trityl chloride yields a Ti0(CO)4 fragment coordinated to an aryl ring of the trityl anion.133,134 Protonation with phenol yielded a dinculear titanium(0) phenolate under H2 evolution.135 Further, [Ti(CO)6]2− forms a complex with azobenzene in which the N]N bond is formally reduced to a dianionic hydrazine ligand, which can be liberated by H2O under formation of [Ti(CO)4(OH)2]2−.136

Scheme 6 Reactivity of group 4 dianionic carbonyl metallates (known only for titanium).

Considerably less is known about the trianionic carbonyl metallates, due to the difficulties for their synthesis as well as their high lability. Reactivity studies on these systems have been analogous to those with their less anionic counterparts, namely derivatisation with alkyl units or organotetrel chlorides, mainly to infer the identity of the trisanionic species (Scheme 7).94,99,105,137–141 Further, protonation give unusual anionic metal hydrides in case of vanadium and group 7 metals, with the metal still in formally negative oxidation states (e.g. [V(CO)5H)]2−, Scheme 7 top).94,99,137–140

Scheme 7 Reactivity of trisanionic carbonyl metallates of group 5 (top), group 7 (middle), and group 9 (bottom).

Very Low Oxidation States in Organometallic Chemistry

93

Tetraanionic carbonyl metallates are known only for group 6. Their synthesis relies on the reduction of the metal(0) precursor (M(CO)4(tmeda) with sodium in liquid ammonia (Scheme 8).

Scheme 8 Reactivity of group 6 tetraanionic carbonyl metallates.

The reactivity of polynuclear carbonyl metallates resembles that of mononuclear metallates (such as hydride formation upon protonation), and can be used for the synthesis of larger homo- and heteronuclear metal carbonylate clusters.75,121–123,142,143 The knowledge on the chemical behaviour of tetraanionic carbonyl metallates is so far limited to derivatisation with organotin chloride or protonation.93,121,144

1.04.2.3

Isonitrile-based metallates

Organoisocyanides are isolobal with carbon monoxide. As such they have been widely used in coordination chemistry as surrogates for CO, with the advantage that organoisocyanides offer the possibility to modulate steric properties and acceptor capabilities via the organic substituent R.145–153 Despite the isolobal analogy, there are significant electronic differences between CO and CNR which stems from the lower electronegativity of the NR group relative to O. This makes organoisocyanides stronger s-donors and weaker p-acids than CO (Fig. 3). As a consequence carbonyls are generally better suited to stabilize metal ions in lower oxidation states, and isonitriles stabilize metal ions in higher oxidation states. Nonetheless, isonitrile metal complexes with highly negatively charged metal ions are known in which pronounced p-backbonding from the metal gives the isonitrile significant carbyne character.154

Fig. 3 Bond interactions between a metal ion and an isonitrile and possible resonance structures.

The first report on an organoisonitrile metallate dates back to 1989 when Cooper presented the synthesis, structure and reactivity of [Co(CNXyl)4]− (Xyl ¼ 2,6-Me2C6H3) which is an analogue of [Co(CO)4]−.155 [Co(CNXyl)4]− was initially synthesized by treatment of K[Co(C2H4)4] with xylyl isocyanide but can also be obtained using [Co(arene)2]− as the cobalt anion source or by reduction of [Co2(CNXyl)8] (Scheme 9).156,157 Similarly, [V(CNXyl)6]− was directly obtained by reduction of the vanadium(0) precursor [V(CNXyl)6] with caesium.151,158 For homoleptic isonitrile metal (di)anions of tantalum, manganese, iron and ruthenium the isonitrile metal mono- (Mn, Nb, Ta) or dihalides (Fe, Ru, Co) were reduced using alkali metal naphthalenides.153,159–162 Mixed carbonyl/isonitrile metallates are known for rhenium (Re(CO)3(CNR)2]−),163 manganese (Mn(CO)3(CNR)2]− and Mn(CO)2(CNR)3]−),164,165 iron and cobalt ([M(CO)2(CNR)2]2− and [M(CO)(CNR)3]2−)166,167 and are mostly obtained by displacing carbon monoxide in metal carbonyl compounds. These compounds can all be seen as analogues of the respective carbonyl metallates, and also obey the 18-electron rule. The use of sterically demanding isonitriles for low-valent complexes was pioneered by Figueroa,168 who reported on unsaturated isonitrile complexes for which the homoleptic carbonyl analogues are elusive, e.g. homoleptic Ni(CNR)3, Co(CNR)4, or heteroleptic Mn(CO)3(CNR)2.154,169–173 The sole example of a homoleptic, formally unsaturated isonitrile metallate is the cobalt(−I) complex ([Co(CNR)3]−).174

94

Very Low Oxidation States in Organometallic Chemistry

Scheme 9 Synthetic routes to homoleptic isonitrile metallates.

As expected from the isoelectronic nature of isonitriles and carbonyls, isonitrile metallates react in a nucleophilic manner. For example, with organo- and organotin halides they give neutral mono- and difunctionalized products (Scheme 10).155,157,161,162,175 Protonation of isonitrile metallates gives metal hydrides whose acidity is lower than their carbonyl analogues, which reflects on the lesser ability of the isonitriles to stabilize a negative charge on the metal.176,177 The reductive capabilities of Mn and Fe complexes were shown by their ability to achieve CO2 reduction, with the iron complex 165,178 yielding reductive disproportionation into CO and CO2− 3 .

Scheme 10 Reactivity of transition metal isonitril metallates.

A remarkable example of the use of isonitrile metallates is the reaction of [Fe(CO)2(CNR)3]2− with BF3 to give the first example of an isolable complex of the BF ligand (Scheme 11). BF is isoelectronic to CO but has little multiple-bond character within the diatomic ligand.167

Scheme 11 Use of a mixed isontrile/carbonyl metallate for isolation of a metal bound fluoroborylene (Tripp ¼ 2,4,6-triisopropylphenyl).

A drawback in the use of isonitriles is the reactivity of their CdN bond in general179 (especially compared to CO) and in low-valent complexes in particular (Scheme 12).174,175,177 In these examples, a substrate can react directly with the isonitrile ligand or insert into the metal-isonitrile bonds.

Scheme 12 Reaction of substrates with the isonitrile ligands in isonitrile metallates.

Very Low Oxidation States in Organometallic Chemistry

1.04.2.4

95

Alkene metallate complexes

Alkenes are a further class of ligands that can stabilize metals in negative oxidation states (forming complexes that here will be designated as alkene metallates). The bonding interaction between an alkene and a metal is generally described by the ChattDewar-Duncanson model.180,181 This model comprises donation of electron density from a p-bond of the alkene into an empty d-orbital of the metal as well as back-donation from a filled d-orbital of the metal into an empty p -orbital of the alkene. The latter is particularly well suited for stabilizing an electron rich metal ion. The synthesis of alkene metallates was initially pursued by Jonas in a quest to obtain binary metal alkene complexes. These studies mainly encompassed compounds of manganese to nickel (Scheme 13).182,183 Thereby Jonas used the approach of displacing cyclopentadienide ligands in metallocenes with lithium or lithium naphthalenide in the presence of alkenes (C2H4, C4H8, cod or COT).184–188 This is accompanied by liberation of lithiated cyclopentadienide (LiCp) as a byproduct.

Scheme 13 Synthesis of known alkene metallates of 3d-transition metals.

For nickel, the nickelate [Ni(cod)2]2− can be easily obtained by reducing the readily available nickel(0) complex Ni(cod)2. The latter is also an intermediate during the reduction of NiCp2 with lithium in the presence of cod. The stability of the alkene metallates also follows the 18-electron rule. This is nicely illustrated for the nickel and the cobalt congeners of M(cod)n2. Whereas in [Co(cod)2]− both ethylene units of the cod ligand bind to the Co−I ion (as in Ni(cod)2), for [Ni(cod)2]2− one cod ligand interacts only through one alkene unit. The manganese compound [Li2(dme)2Mn(cod)2] is a rare case of an open-shell alkene metallate (19 VE). In newer work, alkene complexes of tungsten and tantalum alkene complexes ([WH(C2H4)4]3− and [TaH(Et)(C2H4)3]3−) with formal negative oxidation state were obtained from reaction of metal halides with ethyl lithium (Scheme 14), which acts as a reductant and source of ethylene during the synthesis.189

96

Very Low Oxidation States in Organometallic Chemistry

Scheme 14 Reported alkene metallates of 5d-transition metals.

The reactivity of these alkene metallates was examined as well. With complexes of nickel, cobalt and iron they showed that ethylene or diene ligands can be readily displaced by other alkenes. With cobalt and iron, the reaction with CO gives the expected [Fe(CO)4]2− and [Co(CO)4]−, whereas with nickel no formation of [Ni(CO)3]2− was observed. Remarkably, Jonas also revealed the highly reducing nature of [Ni(cod)2]2− by showing that it is able to reduce anthracene. In [Co(cod)4]− the cyclooctadiene can be protonated or functionalized with nucleophiles such as R2BCl, Me3SiCl, Me3GeCl, or Me3SnCl (Scheme 15). In contrast, for the ethylene complex [Co(C2H4)4]− the electrophile interacts preferably with the metal ion (Scheme 15, bottom).182 It thus shows, that the nucleophilicity of a highly reducing metal can be transferred to a coordinating ligand, which speaks to the partial reduction of the latter.

Scheme 15 Reaction of [Co(Z4-cod)2]− with electrophiles.

In more recent years the alkene metallates of Jonas were reexamined in the context of Kumada-type CdC cross-coupling. Li2(dme)2Fe(C2H4)4 reacts with a silylated allyl bromide to a bisallyl iron(II) compound, in effect behaving as a naked “Fe−II”.190,191 [Ni(cod)2]2− was employed as a molecular precatalyst for CdC cross-coupling192 and served as a precursor for nickel nanoparticles used for olefin hydrogenation.193

1.04.2.4.1

Homoleptic arene-based metallate complexes

The first report on the formation of arene complexes using alkali metal naphthalenides as reductants dates back to 1961. At that time Chatt employed them for the synthesis of zerovalent tris(diphosphine) complexes [M(dmpe)3] of vanadium, chromium and tungsten (dmpe ¼ 1,2-bis(dimethylphosphino ethane; Scheme 16 top), which was subsequently extended to other 3d-metals.194,195 In commemoration of these seminal works, Ellis, who later popularized this approach, proposed to call it the “Chatt reaction.”196 Chatt observed that in some cases the naphthalenide was not just a simple reductant: namely, attempts to isolate the zerovalent tautomer [(Ru(arene)(dmpe)2] led to reversible CdH bond activation to give the hydridoruthenium(II) naphthalenide complex [Ru(H)(naphthyl)(dmpe)2], in which the naphthalenide binds via a s-RudC bond (Scheme 16 bottom).197–199 Although the former could not be directly identified, the reaction of the ruthenium(II) complex with substrates such as CO led to the formation of ruthenium(0) complexes implying the presence of an equilibrium between these two isomers.197

Very Low Oxidation States in Organometallic Chemistry

97

Scheme 16 The “Chatt” reaction (top) and a masked ruthenium(0) complex via reversible CdH bond activation (bottom).

Jonas reported in 1974 that the reaction of the nickel diphosphine or bis(monophosphine) complexes (L2)NiCl (L2 ¼ 1,2-bis(dicyclohexylphosphino)ethane, 1,2-bis(dicyclohexylphosphino)propane or L ¼ tris(cyclohexyl)phosphine) with lithium sand at 0  C in diglyme (bis(2-methoxyethyl) ether) in the presence of aromatic hydrocarbons (e.g. benzene, naphthalene) yielded nickel(0) adducts (L2)Ni(arene).200 The binding mode remained unclear at that point. The lability of the arene substituent was shown via reversible H2 activation that yielded the nickel(II) dihydride (L2)Ni(H)2.201 The structural elucidation of such arene complexes was achieved later with the structurally similar [(dipe)Ni(Z2-naph)] (dipe ¼ 1,2-bis-(diphenylphosphino)ethane).202 The first mention of the synthesis of homoleptic arene complexes using reduced arenes as reductants stemmed from the work of Olivé and Henrici-Olivé.203 They reacted Li(naph) with three equivalents of VCl3 to yield a vanadium(0) bisarene complex [V(naph)2] as revealed by EPR analysis. It was also mentioned that this compound can be reduced further although the resulting product remained elusive. For the analogous reaction with CrCl3 they proposed chromium(I), chromium(−I) and chromium(−III) products but the characterization consisted only of EPR spectra. In a further step, Timms reported in 1983 on elusive anionic bis(arene) complexes of vanadium(I) and titanium(I) by infusion of potassium atoms into a thf solution of the metal halide and arenes at −110  C.204 In a related result, Green a year later described the formation of blue [Ti(C6H6)2]− upon reduction of [Ti(C6H6)2] with a potassium film at ambient temperatures.205 It was then Fochi and Braga who characterized the first anionic arene metallates ([Ti(C6H6)2]− and [Ti(Z6-1,3,5-Me3C6H3)2]−).206 Other titanium(−I) bis(arenes) were reported by Ellis shortly after via the “Chatt reaction,” which he subsequently popularized for a variety of transition metals. One example was the first structurally characterized tris(arene) metallate, namely [Zr(Z4-naph)3]2−.207 This was followed by dianionic examples for the other group 4 metals, as well as monoanionic compounds of niobium and tantalum ([M(arene)3]−).208–210 For vanadium, cobalt and iron bis(arene) compounds [M(arene)2]− can be obtained.156,211–214

1.04.2.4.2

Synthesis and structure of arene metallates

The synthesis of homoleptic anthracenide metal complexes is usually achieved by reacting the respective metal chloride MCln (Scheme 17) with stoichiometric amounts (n+ 1 for odd number elements, n+ 2 for even numbered elements) of sodium/potassium anthracenide or naphthalenide at low temperatures (Table 2). Alkali metal naphthalenides and anthracenides are comparably mild reductants and can be formed in situ upon reacting the arene with the respective alkali metal in donor solvents (mostly THF). They are kinetically stable and even at low temperatures they are soluble in THF or dme. Initial attempts for their use in the reduction of metal salts resulted in the formation of pyrophoric zerovalent metal powders.215 In this capacity they have been used for synthesis by Rieke and others.216,217 To keep the metal in a homogeneous form, suitable ligands must be present during the reduction to kinetically stabilize the metal either in the zerovalent state (as shown above for the synthesis of V(dmpe)3) and/or to allow for further reduction (vide infra). In rare instances also Li2COT was employed as a reductant that leaves a stabilizing p-ligand.218 COT is a rather strong electron acceptor and thus in complexes it is often best described as a persistent dianion.201,212,219 Similar to the carbonyl metallates, the successful isolation and crystallization of these compounds is highly dependent on the complexation of the alkali metal cation by coordinating solvents such as thf or dme, chelating agents (18-crown-6, crypt-2.2.2) or exchange with an organic cation (e.g. NMe+4 or PPN+).

Scheme 17 General synthetic routes for arene metallates.

98

Very Low Oxidation States in Organometallic Chemistry

Table 2

Synthetically accessible arene metallates ( only in situ generated). Complex

Group 4

Group 5

Group 7 Group 8 Group 9

2−

[Ti(arene)3] [Ti(arene)2]− [Zr(arene)3]2− [Hf(arene)3]2− [V(mesH)2]− V(naph)∗ n [Nb(arene)3]− [Nb(COT)3]− [Ta(arene)3]− [Cr(an)2]3−,−∗ [Cr2(arene)2]− [Fe(arene)2]− [Co(arene)2]−

Arene

Reference

naph, anth diphen, mes, naph, C6H6, C6H5Me anth, naph anth mes-H, C6H6 naph anth COT (anth,) naph anth naph anth naph, anth

220 205,206,221 207,220 220 206,222 82 209,210 218 208,210 203 223 211 156,212,213

For nearly all examined transition metals this yields mononuclear, homoleptic compounds (Fig. 4). With chromium, however, the reaction of [Cr(N(SiMe3)2)2(thf )2] with sodium naphthalenide gave a dinuclear metallate with a formal [Cr2]− unit, in which no Cr–Cr interactions were evident. As mentioned before, a mononuclear chromium(−I) and chromium(−III) bis(naphthalenide) complex were also postulated but definite proof is lacking.203 Homoleptic examples of arene metallates of manganese are still missing. The closest examples are [Mn(CO)3]− fragments bound to benzene, which can be used for the acylation of the arene.224–226 A common structural feature of these very low-valent metallate complexes is the bent coordinating arene, often in a Z4-binding mode. This is in contrast to the plethora of arene complexes such as the seminal neutral bis(benzene) chromium complex Cr(Z6-C6H6)2, in which the metal ion lies below the centroid of the planar arene ring.227,228

Fig. 4 Molecular structure of arene metallates.

The [M(arene)] fragment in the arene metallate can thereby be viewed as an arene coordinating to a Mn ion or a radical anion which is more covalently bound to a Mn+1 ion (Fig. 5). The exact electronic structure in most of these homoleptic arene metallates has not yet been addressed in sufficient detail. In a recent report Wolf showed that the Fe-naphthalene interactions are rather covalent. Further, they proposed an admixture of the resonance structures of an Fen(arene)0 with strong p-backbonding and an Fen+1(arene)−, with the latter being the predominant one.

Fig. 5 Bonding modes of arenes in arene metallates.

Very Low Oxidation States in Organometallic Chemistry

99

For the earlier transition metals, the presence of a reduced arene ligand is more likely given their more electropositive nature. Differences in the interaction between the metal and the arene are further implied by unusual Z2- and Z3-arene binding modes found for niobium and tantalum. The situation is complicated by the differences in reducibility of the respective arene and alkene with the qualitative order of naphthalene < anthracene < COT. Calculations on the somewhat related neutral zirconium (0) cyclooctatetraene compound [Zr(Z8-COT)(Zn-COT)] (n ¼ 3, 4) revealed indeed the presence of a zirconium(IV) ion bound to dianionic COT ligands, which might be described as a combination of an allyl and a pentadienyl unit.229 The reduced nature of the arenes can also be used for their direct functionalization with electrophiles.224,225,230 For group 3 metals, lanthanides, and actinides, no homoleptic arene metallates have been isolated so far. However, examples of low-valent group 3 arene complexes (see end of the chapter) give a dianionic character for the metal bound naphthalene or anthracene that can be readily released as neutral molecules.231–234 This overall introduces naphthalenes and larger arenes into the class of redox non-innocent ligands, which can reversibly accept electron density from the metal yet can be easily displaced in their neutral form and thus used as electron reservoir. As a caveat, it has to be mentioned that in some instances the arene might be also displaced as the alkali metal salt if the reduction potential of the reduced metal complex is in the range of the arene radical anion.

1.04.2.4.3

Reactivity

Despite their rather long history the chemical behavior of isolable arene metallates in negative oxidation states is still little developed. The most important reactivity characteristic is the partial or full displacement of the arene for stronger donor ligands. This has been observed for all arene metallates and was employed as a synthetic approach for a variety of carbonyl, isonitrile, alkene or phosphine metallates. The arene can also be displaced by a larger, more easily reduced arene (e.g. anthracene for naphthalene) or alkenes and used for substrate activation. In the following an overview of the work on these arene metallates will be given to give the reader an idea of the state of the art in this field. The group 4 arene metallates were primarily subjected to displacement of the arene by CO or polyenes (Scheme 18).214,221 The most remarkable reaction is the one of Ti(arene)2]2− with white phosphorus. This led to the first carbon-free sandwich complex [Ti(P5)2]2−.235 The reaction of Ti(arene)2]2− with Me3SnCl gave the formally titanium(0) compound [Ti(naph)2(SnMe3)]−.236 In [Ti(naph)2(SnMe3)]− the naphthalene can be subsequently displaced by CO showing the opportunities afforded to introduce ligands in a stepwise fashion.237

Scheme 18 Reactivity of group 4 arene metallates.

The group 5 arene metallates were mainly used in their capacity to undergo complete exchange, e.g. for CO, phosphines or butadiene.208–210,220 It is important to note that the reaction of [V(C6H6)2]− with CO gave not the expected [V(CO)6]− but [V(C6H6)2]0 and an insoluble mixture which was presumed to contain cyclic oligomers of the CO radical anion (e.g. [(CO)4]4−).222,238,239 In case of CO2 several C1 and C2 products were obtained after hydrolysis stemming from the 1e− or 2e− reduction of CO2 (e.g. formate, oxalate).206,222 In contrast, the reaction of the less reducing [V(anth)2]− or [V(anth)2]− with CO or PF3 yielded the homoleptic compounds [V(L)6]− (Scheme 19 top).82,240 For the niobium complex [Nb(anth)3]− the arenes were also fully exchanged using PF3, whereas for P(OMe)3 the substitution remained incomplete in [Nb(anth)2(P(OMe)3)2]− even using an excess of the phosphite (Scheme 19 bottom).

100

Very Low Oxidation States in Organometallic Chemistry

Scheme 19 Synthesis and reactivity of anthracene vanadate (−I) (top) and naphthalene metallates(−I) of niobium and tantalum (bottom).

Butadiene complexes [M(C4H6)3]− of niobium and tantalum, obtained from ligand exchange, can be protonated to give 1-methylallyl-butadiene complexes [M(Z1-MeC3H4)(Z4-C4H6)2].210 This complements observations made for alkene metallates, where the nucleophilicity of the metal ions can be transferred to the bound ligand. For cobalt, homoleptic arene or diene complexes undergo stepwise ligand substitution (Scheme 20).212,213 Similarly, in the heteroleptic [Co(naph)(cod)]− the naphthalenide is displaced first (e.g. reaction with PMe3 gives [Co(cod)(PMe3)2]−).156 Partial displacement of the anthracene can also occur.241

Scheme 20 Ligand exchange reactivity of the cobalt(−I) anthracenide complex (Co(Z4-C14H10)2]−.

The prime example of transferring a naked “Co−I” was the reaction of [Co(naph)2]− with [ArSnCl]2 (Scheme 21) giving [(ArSn)2Co2], a compound with a cyclic [Sn2Co2] core (secondary cobalt-arene interactions with the Ar -ligand are not shown, Ar ¼ C6H3-2,6(C6H2-2,6-iPr)2).242

Scheme 21 Examples of the use of arene cobaltates as source of naked cobalt(−I).

Very Low Oxidation States in Organometallic Chemistry

101

The reaction of “Co−I” with ortho-carborane diphosphetane leads to oxidative addition of the PdP bond and formation of a tetrahedral low-spin cobalt(III) complex.243 In case of the phosphaalkyne P^CtBu substrate dimerization occurs under formation of a diphosphacyclobutadiene complex.244–247 Further, the cobalt(−I) synthon can be reacted with a redox active a-diimine ligands, which are reduced in due process and give a cobalt(+I) complex under two-fold ligand reduction.248,249 Similarly, for an analogous iron(−I) synthon ligand displacement is observed (Scheme 22), which is not always complete even when using a large excess of substrate.211 Iron anthracenide can also dimerize phosphaalkynes, and further trimerize alkynes.244,245,247 Both iron and cobalt metallates can be used as precatalysts for either alkene hydrogenation or CdC cross coupling.250–252

Scheme 22 Reactivity of iron(−I) anthracenide (Fe(Z4-C14H10)2]−.

Together with the related alkene metallates, the presented reactivity examples of arene metallates show the astonishing potential of “naked” metal(−I) ions for a variety of complex formation and bond transformation processes. First, they can be used as precursors for other organometallic metal(−I) complexes via ligand exchange by suitable ligands such as other alkenes/arenes, or P-donor ligands. This can be extended to ligands, that can be either reduced as a result of redox non-innocence or undergo other intraligand bond transformation processes. Second, even more remarkably, substrates can be activated, transformed and incorporated as sole ligands into the metal’s ligand sphere (such as P−5 from P4, or benzene derivatives from alkynes). And last but not least, they can be employed as [M]− for formal halide-metal exchange, a truly astonishing feature.

1.04.2.5

Carbene-based metallates

N-Heterocyclic carbenes (NHC) and cyclic alkyl amino carbenes (cAAC) have found widespread use in the past 10–15 years for the isolation of low-valent main-group compounds as well as low-valent transition metal complexes.253,254 This comes from their strong s-donor character but also significant p-accepting capabilities into the empty p-orbital, especially in case of cAAC’s. In this capacity carbenes were also used successfully in the synthesis of a variety of neutral, formally zerovalent compounds, which are usually found in a linear or otherwise low-coordinate surrounding (Fig. 6).50,255–268

Fig. 6 Examples of two-coordinate, formally zerovalent carbene metal complexes. Half-arrows indicate the (partial) spin-density of one electron divided over both carbene ligands.

This holds also true for bimetallic compounds with ligated [M2]0 units.258,269–273 Closer inspection of the electronic structure of the metal showed that especially for transition metals, the actual zerovalent state of the metal is only found up to a d10 configuration (with the exception of Mn). If the metal ion is reduced further (namely for ZnII, CuI, or AuI (all d10)) the additional electron(s) mostly reside(s) in the empty p-orbital of the carbene carbon.253,257,258 This holds also true for manganese which is not reduced beyond MnI (3d54s1).50 Depending on the further ligand sphere the possible redox-noninnocence of NHC’s and cAAC’s should also be carefully considered for formally monovalent metal compounds (as shown for some two-coordinate mixed amide/NHC ligated cobalt and iron complexes).49 Despite the huge success of stabilising low-valent and even zerovalent metal ions, the further reduction of M(NHC)2,3 or even multinuclear species (such as L! M–M L) is surprisingly unexplored. The only two

102

Very Low Oxidation States in Organometallic Chemistry

systems of a carbene metallate concern the remarkable of reduction of N2-adducts giving compounds bearing M(−I) ions.274,275 They are comparable to phosphine based dinitrogen cobaltates(−I)276 or ferrates(−I).277–279 In case of iron, the labile iron(0) dinitrogen adduct [Fe(cAAC)2N2], obtained by reversible N2 binding to [Fe(cAAC)2] at low temperatures, can be trapped by further reduction to iron(−I) in the presence of 18-crown-6 (Scheme 23).275 The ferrate complex can be silylated winding up in a remarkable end-on silyldiazenido iron(I) unit! These compounds are also capable to catalytically reduce N2 to NH3 (using a proton source) or N(SiMe3)3 (using Me3SiCl) with KC8 as reductant.

Scheme 23 Formation and reactivity of a cAAC based dinitrogen ferrate(−I) complex.

In a similar fashion, the stable cobalt(0) compound [Co(NHC)3N2] can be reduced (by K, Rb, and Cs but not Na; Scheme 24) in the presence of N2 yielding a bis(dinitrogen) cobaltate(−I) complex ().274 Quantum chemical analysis thereby confirmed its formulation as a cobalt(−I) compound with a d10 configuration. Its silylation yielded a cobalt(II) complex bearing a side-on coordinated, dianionic hydrazenide complex via a not fully resolved mechanism. This cobalt system can also be used for catalytic, reductive N2 silylation to N(SiMe3)3.

Scheme 24 Formation and reactivity of a cAAC based dinitrogen cobaltate(−I) complex.

1.04.2.6

Honorable mentions

For lanthanides, solely neutral homoleptic arene complexes are known, as initially shown by Cloke via co-condensation of rare earth metal atoms with arenes (Fig. 7).280–282 This was paralleled by the examination of the reaction of the metal halides of samarium, ytterbium and europium with lithium naphthalenides or sodium anthracenides that gave products of the presumed composition M1–2(naph)∙solventx.283 A consequence of the reducing nature of low-valent lanthanide ions is that the arene ligands appear as mono-, di-, tri or even tetraanions whereas the metal retains its +III oxidation state. In many cases the reducing equivalents come from two lanthanide ions, forming inverse sandwich complexes (Fig. 7).232,284,285

Fig. 7 Examples of rare earth element arene complexes in masked low oxidation states.

As this binding is usually reversible, these arene complexes behave nonetheless as synthons of lower valent lanthanides, which has been used for a rich set of reaction chemistry such as N2 reduction.284,286–292 Organometallic actinide complexes are also found only in positive oxidation states whereas the lowest experimentally verified oxidation state constitutes +II. Analogous compounds are nowadays known for thorium,293,294 uranium,295,296 plutonium297 and neptunium,298 and consist mainly of cyclopentadienide based systems such as [MCp 3]−. Calculations hint at the possibility of

Very Low Oxidation States in Organometallic Chemistry

103

other stable actinide(II) as well as uranium(I) species.299,300 Reversible arene coordination to a low-valent actinide ion, comparable to those of actinides and arene metallates, is restricted to very few examples for thorium and uranium (Fig. 8). For thorium, examples of a thorium(II)233,301 as well as a remarkable thorium(0)234 synthon was reported. The latter is able to cleave N2. In a similar fashion a variety of low-valent uranium arene complexes (serving as synthons for oxidation states as low as UII) have been reported in recent years.302–308

Fig. 8 Examples of actinide element arene complexes in masked low oxidation states (dme ¼ 1,2-dimethoxyethane).

As shown above, the auride anion Au− can be isolated whereas organometallic auride(−I) chemistry is non-existent. However, in recent years it was possible to obtain complexes with a negatively charged gold ion. Haman reported on an ionic gold(−I) complex. The bonding between the gold and boron atoms is described as a single three-center, 2-electron (3c-2e) bond (Fig. 9, left).309 The similarity to B2H6 emphasizes the isolobal relationship between gold and hydrogen.310 Using an anionic aluminyl(I) complex, Aldrich and co-workers isolated a complex bearing an AldAu bond (Fig. 9, right).311 QTAIM analysis showed a strongly polarized Ald+dAud− bond with a calculated charge on gold close to −1. Accordingly, the gold reacts with CO2 in a nucleophilic manner.

Fig. 9 Known auride(−I) complexes as well as their limited reactivity.

1.04.3

Concluding remarks and outlook

The chemistry of organometallics with the metal in (formally) negative oxidation states is a longstanding topic in synthetic inorganic chemistry. To obtain a complex with a metal in a negative oxidation state, the ligands generally need to alleviate the negative charge and high electron density on the metal of through backbonding, which occurs effectively with CO and the isoelectronic isonitriles. The carbonyl metallates are particularly adept for the isolation of complexes bearing a metal in an negative oxidation state of up to −IV, with known examples for nearly all group 4–10 metals. Whereas the chemistry of carbonyl metallates(−I/−II) is relatively developed, accounts on the reactivity of those with even lower oxidation states are scarce, which leaves a wide field for future work. Isolable carbonyl metallates are stable only when they follow the 18-electron rule. Electronically unsaturated analogues can come from the use of isonitrile ligands which are tunable with respect to sterics and electronic properties. As a result, the number of metallate complexes bearing sterically demanding isonitriles is still fairly limited but the initial studies are highly promising. A further approach is the use of alkenes that stabilize the metal anion by p-backbonding. As the alkenes can be fully displaced, respective metallates can be seen as masked form of a metal anion. Homoleptic arene metallates (group 4, 5 and 7–9) can also be seen in a similar way. In these, the coordinating arene can reversibly store an electron as a metal bound arene radical (di)anion. The capacity of arene or alkene stabilized “naked” metal anions is less developed but offer great promise for the synthesis other low-molecular metal containing compounds. As such they can mediate a variety of different reactions, such as ligand

104

Very Low Oxidation States in Organometallic Chemistry

substitution, coordination of ligands in conjunction with ligand reduction or even intra-ligand bond cleavage, and formation of new bonds or even new ligands in case of elements as substrates! The carbene based metallate chemistry is much less explored, but given the propensity of N-heterocyclic and cyclic alkyl amino carbenes to stabilize low and zero-valent metal ions, and the ability to provide steric protection with these ligands, future studies are expected to yield continued remarkable outcomes. Metallates of the heavier transition metals are scarce and for lanthanides and actinides are missing, which gives a further “playground” for future endeavors. Overall, masked metal anions in form of homoleptic arene metallates are potentially capable of linking the chemistry in solution with that under more extreme conditions such as in gas-phase or in a low-temperature matrix. Further, they may give deeper, fundamental insights in terms of metal-substrate bonding and may pave the way to novel bond activation modes.

Acknowledgement The author thanks the Deutsche Forschungsgesellschaft (DFG grant WE 5627/4–1) and the Philipps-University for financial support.

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. 29. 30. 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43. 44. 45. 46. 47. 48. 49.

Thanh, N. T. K.; Maclean, N.; Mahiddine, S. Chem. Rev. 2014, 114, 7610–7630. Joannis, M. C. R. Hebd. Seances Acad. Sci. 1891, 113, 795. Hieber, W.; Leutert, F. Z. Anorg. Allg. Chem. 1932, 204, 145–164. Smyth, F. H. J. Am. Chem. Soc. 1917, 39, 1299–1312. Zintl, E.; Goubean, J.; Dullenkopf, W. Z. Phys. Chem. 1931, 154A, 1–46. Kraus, C. A. J. Am. Chem. Soc. 1922, 44, 1216–1239. Peck, E. B. J. Am. Chem. Soc. 1918, 40, 335–347. Kraus, C. A. Trans. Am. Electrochem. Soc. 1924, 45, 175. Zintl, E. Angew. Chem. 1939, 52, 1–6. Corbett, J. D. Chem. Rev. 1985, 85, 383–397. Mingos, D. M. P., Corbett, J. D., Eds.; In Structural and Electronic Paradigms in Cluster Chemistry; Structure and Bonding; Springer-Verlag: Berlin, 1997; vol. 87. Sevov, S. C.; Goicoechea, J. M. Organometallics 2006, 25, 5678–5692. Scharfe, S.; Kraus, F.; Stegmaier, S.; Schier, A.; Fässler, T. F. Angew. Chem. Int. Ed. 2011, 50, 3630–3670. Wilson, R. J.; Weinert, B.; Dehnen, S. Dalton Trans. 2018, 47, 14861–14869. Wilson, R. J.; Lichtenberger, N.; Weinert, B.; Dehnen, S. Chem. Rev. 2019, 119, 8506–8554. Sloane, R. H.; Press, R. Proc. R. Soc. Lond. A 1938, 168, 284–301. Arnot, F. L.; Beckett, C. Proc. R. Soc. Lond. A 1938, 168, 103–122. Arnot, F. L.; Milligan, J. C. Proc. R. Soc. Lond. A 1936, 156, 538–560. Skinner, H. A.; Pritchard, H. O. Trans. Faraday Soc. 1953, 49, 1254. Zollweg, R. J. J. Chem. Phys. 1969, 50, 4251–4261. Rienstra-Kiracofe, J. C.; Tschumper, G. S.; Schaefer, H. F.; Nandi, S.; Ellison, G. B. Chem. Rev. 2002, 102, 231–282. Sloane, R. H.; Love, H. M. Nature 1947, 159, 302–303. Kalcher, J.; Sax, A. F. Chem. Rev. 1994, 94, 2291–2318. Salvador, P.; Vos, E.; Corral, I.; Andrada, D. M. Angew. Chem. Int. Ed. 2021, 60, 1498. Liu, G.; Fedik, N.; Martinez-Martinez, C.; Ciborowski, S. M.; Zhang, X.; Boldyrev, A. I.; Bowen, K. H. Angew. Chem. 2019, 131, 13927–13931. Pan, S.; Frenking, G. Angew. Chem. Int. Ed. 2020, 59, 8756–8759. Foroutan-Nejad, C. Angew. Chem. Int. Ed. 2020, 59, 20900–20903. Pyper, N. C.; Edwards, P. P. J. Am. Chem. Soc. 2000, 122, 5092–5099. Edwards, P. P.; Ellaboudy, A. S.; Holton, D. M. Nature 1985, 317, 242–244. Jansen, M. Chem. Soc. Rev. 2008, 37, 1826–1835. Tehan, F. J.; Barnett, B. L.; Dye, J. L. J. Am. Chem. Soc. 1974, 96, 7203–7208. Dye, J. L. J. Phys. Chem. 1984, 88, 3842–3846. Karpov, A.; Nuss, J.; Wedig, U.; Jansen, M. Angew. Chem. Int. Ed. 2003, 42, 4818–4821. Karpov, A.; Wedig, U.; Dinnebier, R. E.; Jansen, M. Angew. Chem. Int. Ed. 2005, 44, 770–773. Maher, J. M.; Fox, J. R.; Foxman, B. M.; Cooper, N. J. J. Am. Chem. Soc. 1984, 106, 2347–2353. Rosenthal, U. Angew. Chem. Int. Ed. 2003, 42, 1794–1798. Rosenthal, U.; Pellny, P. M.; Kirchbauer, F. G.; Burlakov, V. V. Acc. Chem. Res. 2000, 33, 119–129. Buchwald, S. L.; Nielsen, R. B. Chem. Rev. 1988, 88, 1047–1058. Wink, D. J.; Fox, J. R.; Cooper, N. J. J. Am. Chem. Soc. 1985, 107, 5012–5014. Jørgensen, C. Coord. Chem. Rev. 1966, 1, 164–178. Bowman, A. C.; England, J.; Sproules, S.; Weyhermüller, T.; Wieghardt, K. Inorg. Chem. 2013, 52, 2242–2256. Mézailles, N.; Rosa, P.; Ricard, L.; Mathey, F.; Le Floch, P. Organometallics 2000, 19, 2941–2943. Enemark, J. H.; Feltham, R. D. Coord. Chem. Rev. 1974, 13, 339–406. Munz, D. Organometallics 2018, 37, 275–289. Dzik, W. I.; Zhang, X. P.; de Bruin, B. Inorg. Chem. 2011, 50, 9896–9903. Caulton, K. G. Eur. J. Inorg. Chem. 2012, 2012, 435–443. Lyaskovskyy, V.; de Bruin, B. ACS Catal. 2012, 2, 270–279. Flisak, Z.; Sun, W.-H. ACS Catal. 2015, 5, 4713–4724. Danopoulos, A. A.; Braunstein, P.; Monakhov, K. Y.; van Leusen, J.; Kögerler, P.; Clémancey, M.; Latour, J.-M.; Benayad, A.; Tromp, M.; Rezabal, E.; Frison, G. Dalton Trans. 2017, 46, 1163–1171.

Very Low Oxidation States in Organometallic Chemistry

105

50. Samuel, P. P.; Mondal, K. C.; Roesky, H. W.; Hermann, M.; Frenking, G.; Demeshko, S.; Meyer, F.; Stückl, A. C.; Christian, J. H.; Dalal, N. S.; Ungur, L.; Chibotaru, L. F.; Pröpper, K.; Meents, A.; Dittrich, B. Angew. Chem. 2013, 125, 12033–12037. 51. Knijnenburg, Q.; Gambarotta, S.; Budzelaar, P. H. M. Dalton Trans. 2006, 5442–5448. 52. Kubin, M.; Guo, M.; Kroll, T.; Löchel, H.; Källman, E.; Baker, M. L.; Mitzner, R.; Gul, S.; Kern, J.; Föhlisch, A.; Erko, A.; Bergmann, U.; Yachandra, V.; Yano, J.; Lundberg, M.; Wernet, P. Chem. Sci. 2018, 9, 6813–6829. 53. George, S. D.; Brant, P.; Solomon, E. I. J. Am. Chem. Soc. 2005, 127, 667–674. 54. Gao, C.; Macetti, G.; Overgaard, J. Inorg. Chem. 2019, 58, 2133–2139. 55. Nandi, A.; Kozuch, S. Chem. Eur. J. 2020, 26, 759–772. 56. Jerabek, P.; Schwerdtfeger, P.; Frenking, G. J. Comput. Chem. 2019, 40, 247–264. 57. Walroth, R. C.; Lukens, J. T.; MacMillan, S. N.; Finkelstein, K. D.; Lancaster, K. M. J. Am. Chem. Soc. 2016, 138, 1922–1931. 58. Jiang, W.; DeYonker, N. J.; Wilson, A. K. J. Chem. Theory Comput. 2012, 8, 460–468. 59. Escudero, D.; Thiel, W. J. Chem. Phys. 2014, 140. 194105-1-8. 60. Husch, T.; Freitag, L.; Reiher, M. J. Chem. Theory Comput. 2018, 14, 2456–2468. 61. Hoffmann, R.; Alvarez, S.; Mealli, C.; Falceto, A.; Cahill, T. J.; Zeng, T.; Manca, G. Chem. Rev. 2016, 116, 8173–8192. 62. Carsch, K. M.; DiMucci, I. M.; Iovan, D. A.; Li, A.; Zheng, S.-L.; Titus, C. J.; Lee, S. J.; Irwin, K. D.; Nordlund, D.; Lancaster, K. M.; Betley, T. A. Science 2019, 365, 1138–1143. 63. Steen, J. S.; Knizia, G.; Klein, J. E. M. N. Angew. Chem. Int. Ed. 2019, 58, 13133–13139. 64. Snyder, J. P. Angew. Chem. Int. Ed. 1995, 34, 80–81. 65. Poli, R. Chem. Rev. 1996, 96, 2135–2204. 66. Müller, I.; Munz, D.; Werncke, C. G. Inorg. Chem. 2020, 59, 9521–9537. 67. Dapprich, S.; Frenking, G. J. Phys. Chem. 1995, 99, 9352–9362. 68. Szilagyi, R. K.; Frenking, G. Organometallics 1997, 16, 4807–4815. 69. Ehlers, A. W.; Dapprich, S.; Vyboishchikov, S. F.; Frenking, G. Organometallics 1996, 15, 105–117. 70. Hock, H.; Stuhlmann, H. Ber. dtsch. Chem. Ges. A/B 1929, 62, 2690–2693. 71. Hieber, W.; Leutert, F. Naturwissenschaften 1931, 19, 360–361. 72. Krumholz, P.; Stettiner, H. M. A. J. Am. Chem. Soc. 1949, 71, 3035–3039. 73. Dewar, J.; Jones, H. O. Proc. R. Soc. Lond. A 1905, 76, 558–577. 74. Case, J. R.; Whiting, M. C. J. Chem. Soc. 1960, 4632. 75. Hieber, W.; Brendel, G. Z. Anorg. Allg. Chem. 1957, 289, 324–337. 76. Hieber, W.; Abeck, W.; Sedlmeier, J. Angew. Chem. 1952, 64, 480. 77. Hieber, W.; Wagner, G. Z. Naturforsch. B Chem. Sci. 1957, 12, 478–479. 78. Edgell, W. F.; Lyford, J. Inorg. Chem. 1970, 9, 1932–1933. 79. Hieber, W.; Beck, W.; Braun, G. Angew. Chem. 1960, 72, 795–801. 80. Hieber, W.; Abeck, W.; Platzer, H. K. Z. anorg. allg. Chem. 1955, 280, 241–251. 81. King, R. B. Reactions of alkali metal derivatives of metal carbonyls and related compounds. In Advances in Organometallic Chemistry; Stone, F. G. A., West, R., Eds.; Academic Press: New York, London, 1964;; pp 157–256. 82. Barybin, M. V.; Pomije, M. K.; Ellis, J. E. Inorg. Chim. Acta 1998, 269, 58–62. 83. Ellis, J. E.; Warnock, G. F.; Barybin, M. V.; Pomije, M. K. Chem. Eur. J. 1995, 1, 521–527. 84. Dewey, C. G.; Ellis, J. E.; Fjare, K. L.; Pfahl, K. M.; Warnock, G. F. P. Organometallics 1983, 2, 388–391. 85. Datta, S.; Wreford, S. S. Inorg. Chem. 1977, 16, 1134–1137. 86. Ellis, J. E. Organometallics 2003, 22, 3322–3338. 87. Ellis, J. E.; Chi, K. M. J. Am. Chem. Soc. 1990, 112, 6022–6025. 88. Wilson, R. D.; Bau, R. J. Am. Chem. Soc. 1974, 96, 7601–7602. 89. Ellis, J. E.; Fjare, K. L.; Hay, T. G. J. Am. Chem. Soc. 1981, 103, 6100–6106. 90. Ellis, J. E.; Hagen, G. P. J. Am. Chem. Soc. 1974, 96, 7825–7826. 91. H. Behrens, Adv. Organomet. Chem. 18, 1980, 1–53. 92. Ellis, J. E.; Hentges, S. G.; Kalina, D. G.; Hagen, G. P. J. Organomet. Chem. 1975, 97, 79–93. 93. Ellis, J. E.; Parnell, C. P.; Hagen, G. P. J. Am. Chem. Soc. 1978, 100, 3605–3607. 94. Ellis, J. E.; Faltynek, R. A. J. Chem. Soc., Chem. Commun. 1975, 966. 95. Hieber, W.; Mühlbauer, F.; Ehmann, E. A. Ber. dtsch. Chem. Ges. A/B 1932, 65, 1090–1101. 96. Ellis, J. E.; Barger, P. T.; Winzenburg, M. L. J. Chem. Soc., Chem. Commun. 1977, 686–687. 97. Calderazzo, F.; Englert, U.; Pampaloni, G.; Pelizzi, G.; Zamboni, R. Inorg. Chem. 1983, 22, 1865–1870. 98. Barybin, M. V.; Ellis, J. E.; Pomije, M. K.; Tinkham, M. L.; Warnock, G. F. Inorg. Chem. 1998, 37, 6518–6527. 99. Warnock, G. F. P.; Sprague, J.; Fjare, K. L.; Ellis, J. E. J. Am. Chem. Soc. 1983, 105, 672. 100. Hileman, J. C.; Huggins, D. K.; Kaesz, H. D. Inorg. Chem. 1962, 1, 933–938. 101. Cotton, J. D.; Bruce, M. I.; Stone, F. G. A. J. Chem. Soc. A 1968, 2162. 102. Chini, P.; Martinengo, S. Inorg. Chim. Acta 1969, 3, 21–24. 103. Hieber, W.; Braun, G. Z. Naturforsch. B Chem. Sci. 1959, 14, 132–133. 104. L’Eplattenier, F.; Pélichet, M. C. HCA 1970, 53, 1091–1099. 105. Allen, J. M.; Brennessel, W. W.; Buss, C. E.; Ellis, J. E.; Minyaev, M. E.; Pink, M.; Warnock, G. F.; Winzenburg, M. L.; Young, V. G. Inorg. Chem. 2001, 40, 5279–5284. 106. Chi, C.; Pan, S.; Meng, L.; Luo, M.; Zhao, L.; Zhou, M.; Frenking, G. Angew. Chem. Int. Ed. 2019, 58, 1732–1738. 107. Jin, J.; Yang, T.; Xin, K.; Wang, G.; Jin, X.; Zhou, M.; Frenking, G. Angew. Chem. Int. Ed. 2018, 57, 6236–6241. 108. Squires, R. R. Chem. Rev. 1987, 87, 623–646. 109. Pike, R. D. Organometallics 2012, 31, 7647–7660. 110. Strauss, S. H. J. Chem. Soc., Dalton Trans. 2000, 1–6. 111. Lorenz, C.; Kaas, M.; Korber, N. Z. Anorg. Allg. Chem. 2018, 644, 1678–1680. 112. Blanchard, A. A. Chem. Rev. 1940, 26, 409–422. 113. Chin, H. B.; Bau, R. J. Am. Chem. Soc. 1976, 98, 2434–2439. 114. Klüfers, P. Z. Kristallogr. Cryst. Mater. 1984, 167, 275–286. 115. Bistoni, G.; Rampino, S.; Scafuri, N.; Ciancaleoni, G.; Zuccaccia, D.; Belpassi, L.; Tarantelli, F. Chem. Sci. 2016, 7, 1174–1184. 116. Edgell, W. F.; Watts, A. T.; Lyford, J.; Risen, W. M. J. Am. Chem. Soc. 1966, 88, 1815. 117. M.Y. Darensbourg, Prog. Inorg. Chem. 33, 221–274. 118. Stammreich, H.; Kawai, K.; Tavares, Y.; Krumholz, P.; Behmoiras, J.; Bril, S. J. Chem. Phys. 1960, 32, 1482–1487. 119. Edgell, W. F.; Yang, M. T.; Koizumi, N. J. Am. Chem. Soc. 1965, 87, 2563–2567.

106

120. 121. 122. 123. 124. 125. 126. 127. 128. 129. 130. 131. 132. 133. 134. 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. 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. 191.

Very Low Oxidation States in Organometallic Chemistry

Geier, J.; Willner, H.; Lehmann, C. W.; Aubke, F. Inorg. Chem. 2007, 46, 7210–7214. Lin, J. T.; Hagen, G. P.; Ellis, J. E. J. Am. Chem. Soc. 1983, 105, 2296–2303. Brunet, J.-J.; Neibecker, D.; Shyam Srivastava, R. J. Organomet. Chem. 1993, 461, 169–172. Nagel, C. C.; Bricker, J. C.; Alway, D. G.; Shore, S. G. J. Organomet. Chem. 1981, 219, C9–C12. Behrens, H.; Lohöfer, F. Ber. Dtsch. Chem. Ges. 1961, 94, 1391–1402. Hoffmann, R. Angew. Chem. Int. Ed. 1982, 21, 711–724. Collman, J. P. Acc. Chem. Res. 1975, 8, 342–347. Ellis, J. E. J. Organomet. Chem. 1975, 86, 1–56. Berti, B.; Bortoluzzi, M.; Cesari, C.; Femoni, C.; Iapalucci, M. C.; Mazzoni, R.; Vacca, F.; Zacchini, S. Eur. J. Inorg. Chem. 2019, 2019, 3084–3093. Collman, J. P.; Vastine, F. D.; Roper, W. R. J. Am. Chem. Soc. 1968, 90, 2282–2287. Braunschweig, H.; Dewhurst, R. D.; Schneider, A. Chem. Rev. 2010, 110, 3924–3957. Cowley, A. H.; Lomelí, V.; Voigt, A. J. Am. Chem. Soc. 1998, 120, 6401–6402. Chi, K. M.; Frerichs, S. R.; Philson, S. B.; Ellis, J. E. J. Am. Chem. Soc. 1988, 110, 303–304. Fischer, P. J.; Ahrendt, K. A.; Young, V. G.; Ellis, J. E. Organometallics 1998, 17, 13–15. Fischer, P. J.; Young, V. G., Jr.; Ellis, J. E. Chem. Commun. 1997, 1249–1250. Fischer, P. J.; Yuen, P.; Young, V. G.; Ellis, J. E. J. Am. Chem. Soc. 1997, 119, 5980–5981. Tripepi, G.; Young, V. G.; Ellis, J. E. J. Organomet. Chem. 2000, 593-594, 354–360. Warnock, G. F. P.; Moodie, L. C.; Ellis, J. E. J. Am. Chem. Soc. 1989, 111, 2131–2141. Beck, W. Angew. Chem. Int. Ed. 1991, 30, 168–169. Ellis, J. E.; Hayes, T. G.; Stevens, R. E. J. Organomet. Chem. 1981, 216, 191–209. Warnock, G. F. P.; Moodie, L. C.; Ellis, J. E. ChemInform 1989, 20. Ellis, J. E.; Barger, P. T.; Winzenburg, M. L.; Warnock, G. F. J. Organomet. Chem. 1990, 383, 521–530. Corraine, M. S.; Atwood, J. D. Organometallics 1991, 10, 2647–2651. Krause, J. A.; Siriwardane, U.; Salupo, T. A.; Wermer, J. R.; Knoeppel, D. W.; Shore, S. G. J. Organomet. Chem. 1993, 454, 263–271. Rochfort, G. L.; Ellis, J. E. J. Organomet. Chem. 1983, 250, 265–276. Carpenter, A. E.; Mokhtarzadeh, C. C.; Ripatti, D. S.; Havrylyuk, I.; Kamezawa, R.; Moore, C. E.; Rheingold, A. L.; Figueroa, J. S. Inorg. Chem. 2015, 54, 2936–2944. Cotton, F. A.; Zingales, F. J. Am. Chem. Soc. 1961, 83, 351–355. Sarapu, A. C.; Fenske, R. F. Inorg. Chem. 1972, 11, 3021–3025. P.M. Treichel, Adv. Organomet. Chem. 11,1973, 21–86. Yamamoto, Y. Coord. Chem. Rev. 1980, 32, 193–233. Weber, L. Angew. Chem. Int. Ed. 1998, 37, 1515–1517. Barybin, M. V.; Young, V. G.; Ellis, J. E. J. Am. Chem. Soc. 2000, 122, 4678–4691. Barybin, M. V.; Meyers, J. J.; Neal, B. M. Renaissance of isocyanoarenes as ligands in low-valent organometallics. In Isocyanide Chemistry; Nenajdenko, V., Ed.; Wiley-VCH: Weinheim, Germany, 2012;; pp 493–529. Barybin, M. V.; Brennessel, W. W.; Kucera, B. E.; Minyaev, M. E.; Sussman, V. J.; Young, V. G.; Ellis, J. E. J. Am. Chem. Soc. 2007, 129, 1141–1150. Mokhtarzadeh, C. C.; Moore, C. E.; Rheingold, A. L.; Figueroa, J. S. J. Am. Chem. Soc. 2018, 140, 8100–8104. Warnock, G. F.; Cooper, N. J. Organometallics 1989, 8, 1826–1827. Brennessel, W. W.; Ellis, J. E. Inorg. Chem. 2012, 51, 9076–9094. Leach, P. A.; Geib, S. J.; Corella, J. A.; Warnock, G. F.; Cooper, N. J. J. Am. Chem. Soc. 1994, 116, 8566–8574. Barybin, M. V.; Young, V. G.; Ellis, J. E. J. Am. Chem. Soc. 1998, 120, 429–430. Brennessel, W. W.; Ellis, J. E. Angew. Chem. Int. Ed. 2007, 46, 598–600. Barybin, M. V.; Young, V. G.; Ellis, J. E. J. Am. Chem. Soc. 1999, 121, 9237–9238. Utz, T. L.; Leach, P. A.; Geib, S. J.; John Cooper, N. N. Chem. Commun. 1997, 847–848. Corella, J. A.; Thompson, R. L.; Cooper, N. J. Angew. Chem. Int. Ed. 1992, 31, 83–84. Salsi, F.; Neville, M.; Drance, M.; Hagenbach, A.; Chan, C.; Figueroa, J. S.; Abram, U. Chem. Commun. 2020, 56, 7009–7012. Stewart, M. A.; Moore, C. E.; Ditri, T. B.; Labios, L. A.; Rheingold, A. L.; Figueroa, J. S. Chem. Commun. 2011, 47, 406–408. Agnew, D. W.; Sampson, M. D.; Moore, C. E.; Rheingold, A. L.; Kubiak, C. P.; Figueroa, J. S. Inorg. Chem. 2016, 55, 12400–12408. Carpenter, A. E.; Chan, C.; Rheingold, A. L.; Figueroa, J. S. Organometallics 2016, 35, 2319–2326. Drance, M. J.; Sears, J. D.; Mrse, A. M.; Moore, C. E.; Rheingold, A. L.; Neidig, M. L.; Figueroa, J. S. Science 2019, 363, 1203–1205. Ditri, T. B.; Fox, B. J.; Moore, C. E.; Rheingold, A. L.; Figueroa, J. S. Inorg. Chem. 2009, 48, 8362–8375. Margulieux, G. W.; Weidemann, N.; Lacy, D. C.; Moore, C. E.; Rheingold, A. L.; Figueroa, J. S. J. Am. Chem. Soc. 2010, 132, 5033–5035. Barnett, B. R.; Figueroa, J. S. Chem. Commun. 2016, 52, 13829–13839. Emerich, B. M.; Moore, C. E.; Fox, B. J.; Rheingold, A. L.; Figueroa, J. S. Organometallics 2011, 30, 2598–2608. Agnew, D. W.; Moore, C. E.; Rheingold, A. L.; Figueroa, J. S. Angew. Chem. Int. Ed. 2015, 54, 12673–12677. Fox, B. J.; Millard, M. D.; DiPasquale, A. G.; Rheingold, A. L.; Figueroa, J. S. Angew. Chem. Int. Ed. 2009, 48, 3473–3477. Carpenter, A. E.; Margulieux, G. W.; Millard, M. D.; Moore, C. E.; Weidemann, N.; Rheingold, A. L.; Figueroa, J. S. Angew. Chem. Int. Ed. 2012, 51, 9412–9416. Utz, T. L.; Leach, P. A.; Geib, S. J.; Cooper, N. J. Organometallics 1997, 16, 4109–4114. Mokhtarzadeh, C. C.; Margulieux, G. W.; Carpenter, A. E.; Weidemann, N.; Moore, C. E.; Rheingold, A. L.; Figueroa, J. S. Inorg. Chem. 2015, 54, 5579–5587. Carpenter, A. E.; Rheingold, A. L.; Figueroa, J. S. Organometallics 2016, 35, 2309–2318. Mokhtarzadeh, C. C.; Moore, C. E.; Rheingold, A. L.; Figueroa, J. S. Angew. Chem. Int. Ed. 2017, 56, 10894–10899. Qiu, G.; Ding, Q.; Wu, J. Chem. Soc. Rev. 2013, 42, 5257–5269. Chatt, J.; Duncanson, L. A. J. Chem. Soc. 1953, 2939. Dewar, M. Bull. Soc. Chem. Fr. 1951, 18, C79. Jonas, K. Angew. Chem. Int. Ed. 1985, 24, 295–311. Fischer, K.; Jonas, K.; Wilke, G. Angew. Chem. Int. Ed. 1973, 12, 565–566. Jonas, K.; Häselhoff, C.-C.; Goddard, R.; Krüger, C. Inorg. Chim. Acta 1992, 198-200, 533–541. Jonas, K. Angew. Chem. Int. Ed. 1975, 14, 752–753. Jonas, K.; Mynott, R.; Krüger, C.; Sekutowski, J. C.; Tsay, Y.-H. Angew. Chem. Int. Ed. 1976, 15, 767–768. Jonas, K.; Schieferstein, L.; Krüger, C.; Tsay, Y.-H. Angew. Chem. Int. Ed. 1979, 18, 550–551. Jonas, K.; Schieferstein, L. Angew. Chem. Int. Ed. 1979, 18, 549–550. Morse, P. M.; Shelby, Q. D.; Kim, D. Y.; Girolami, G. S. Organometallics 2008, 27, 984–993. Smith, J. D.; Hanusa, T. P.; Young, V. G. J. Am. Chem. Soc. 2001, 123, 6455–6456. Fürstner, A.; Martin, R.; Krause, H.; Seidel, G.; Goddard, R.; Lehmann, C. W. J. Am. Chem. Soc. 2008, 130, 8773–8787.

Very Low Oxidation States in Organometallic Chemistry

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. 233. 234. 235. 236. 237. 238. 239. 240. 241. 242. 243. 244. 245. 246. 247. 248. 249. 250. 251. 252. 253. 254. 255. 256. 257. 258. 259. 260. 261.

Nattmann, L.; Lutz, S.; Ortsack, P.; Goddard, R.; Cornella, J. J. Am. Chem. Soc. 2018, 140, 13628–13633. Maier, T. M.; Sandl, S.; Melzl, P.; Zweck, J.; Jacobi von Wangelin, A.; Wolf, R. Chem. Eur. J. 2020, 26, 6113–6117. Chatt, J.; Watson, H. R. Nature 1961, 189, 1003–1004. Chatt, J.; Watson, H. R. J. Chem. Soc. 1962, 2545–2549. Ellis, J. E. Dalton Trans. 2019, 48, 9538–9563. Chatt, J.; Davidson, J. M. J. Chem. Soc. 1965, 843. Ibekwe, S. D.; Kilbourn, B. T.; Raeburn, U. A.; Russell, D. R. J. Chem. Soc. D 1969, 433. Gregory, U. A.; Ibekwe, S. D.; Kilbourn, B. T.; Russell, D. R. J. Chem. Soc. A 1971, 1118. Jonas, K. J. Organomet. Chem. 1974, 78, 273–279. Jonas, K.; Krüger, C. Angew. Chem. Int. Ed. 1980, 19, 520–537. Scott, F.; Krüger, C.; Betz, P. J. Organomet. Chem. 1990, 387, 113–121. Henrici-Olive, G.; Olive, S. J. Am. Chem. Soc. 1970, 92, 4831–4834. Hawker, P. N.; Timms, P. L. J. Chem. Soc., Dalton Trans. 1983, 1123. Bandy, J. A.; Berry, A.; Green, M. L. H.; Perutz, R. N.; Prout, K.; Verpeaux, J.-N. J. Chem. Soc., Chem. Commun. 1984, 729–731. Fochi, G.; Braga, D.; Sabatino, P. Organometallics 1988, 7, 565–566. Jang, M.; Ellis, J. E. Angew. Chem. Int. Ed. 1994, 33, 1973–1975. Brennessel, W. W.; Ellis, J. E.; Pomije, M. K.; Sussman, V. J.; Urnezius, E.; Young, V. G. J. Am. Chem. Soc. 2002, 124, 10258–10259. Brennessel, W. W.; Ellis, J. E.; Roush, S. N.; Strandberg, B. R.; Woisetschläger, O. E.; Young, V. G. Chem. Commun. 2002, 2356–2357. Sussman, V. J.; Ellis, J. E. Angew. Chem. Int. Ed. 2008, 47, 484–489. Brennessel, W. W.; Jilek, R. E.; Ellis, J. E. Angew. Chem. Int. Ed. 2007, 46, 6132–6136. Brennessel, W. W.; Young, J. V. G.; Ellis, J. E. Angew. Chem. Int. Ed. 2002, 41, 1211–1215. Brennessel, W. W.; Young, V. G.; Ellis, J. E. Angew. Chem. Int. Ed. 2006, 45, 7268–7271. Kucera, B. E.; Jilek, R. E.; Brennessel, W. W.; Ellis, J. E. Acta Crystallogr. C 2014, 70, 749–753. Scott, N. D.; Walker, J. F.; Hansley, V. L. Sodium Naphthalene. I. A New Method for the Preparation of Addition Compounds of Alkali Metals and Polycyclic Aromatic Hydrocarbons; E. I. du Pont de Nemours and Co., Ed., 1936 Rieke, R. D.; Li, P. T.-J.; Burns, T. P.; Uhm, S. T. J. Org. Chem. 1981, 46, 4323–4324. Li Chu, T.; Friel, J. V. J. Am. Chem. Soc. 1955, 77, 5838–5840. Guggenberger, L. J.; Schrock, R. R. J. Am. Chem. Soc. 1975, 97, 6693–6700. Edelmann, F. T. New J. Chem. 2011, 35, 517–528. Jilek, R. E.; Jang, M.; Smolensky, E. D.; Britton, J. D.; Ellis, J. E. Angew. Chem. 2008, 120, 8820–8823. Blackburn, D. W.; Britton, D.; Ellis, J. E. Angew. Chem. Int. Ed. 1992, 31, 1495–1498. Fochi, G. J. Organomet. Chem. 1988, 350, C1–C3. Labrum, N. S.; Losovyj, Y.; Caulton, K. G. Chem. Commun. 2018, 54, 12397–12399. Shao, L.; Badger, P. D.; Geib, S. J.; Cooper, N. J. Organometallics 2004, 23, 5939–5943. Lee, S.; Geib, S. J.; Cooper, N. J. J. Am. Chem. Soc. 1995, 117, 9572–9573. Reingold, J. A.; Virkaitis, K. L.; Carpenter, G. B.; Sun, S.; Sweigart, D. A.; Czech, P. T.; Overly, K. R. J. Am. Chem. Soc. 2005, 127, 11146–11158. Pampaloni, G. Coord. Chem. Rev. 2010, 254, 402–419. Fischer, E. O.; Hafner, W. Z. Naturforsch. B Chem. Sci. 1955, 10, 665–668. Strauch, H. C.; Bergander, K.; Kehr, G.; Fröhlich, R.; Erker, G. Ber. Dtsch. Chem. Ges. 1999, 1999, 1461–1466. Veauthier, J. M.; Chow, A.; Fraenkel, G.; Geib, S. J.; Cooper, N. J. Organometallics 2000, 19, 661–671. Huang, W.; Khan, S. I.; Diaconescu, P. L. J. Am. Chem. Soc. 2011, 133, 10410–10413. Fryzuk, M. D.; Jafarpour, L.; Kerton, F. M.; Love, J. B.; Rettig, S. J. Angew. Chem. Int. Ed. 2000, 39, 767–770. Korobkov, I.; Gambarotta, S.; Yap, G. P. A. Angew. Chem. Int. Ed. 2003, 42, 814–818. Korobkov, I.; Gambarotta, S.; Yap, G. P. A. Angew. Chem. Int. Ed. 2003, 42, 4958–4961. Urnius, E.; Brennessel, W. W.; Cramer, C. J.; Ellis, J. E.; Schleyer, P. V. R. Science 2002, 295, 832–834. Ellis, J. E.; Blackburn, D. W.; Yuen, P.; Jang, M. J. Am. Chem. Soc. 1993, 115, 11616–11617. Ellis, J. E.; Yuen, P.; Jang, M. J. Organomet. Chem. 1996, 507, 283–286. Fochi, G.; Runjuan, X.; Colligiani, A. J. Chem. Soc., Dalton Trans. 1990, 2551. Lednor, P. W.; Versloot, P. C. J. Chem. Soc., Chem. Commun. 1983, 284. Schmidt, H.; Rehder, D. Transition. Met. Chem. 1980, 5, 214–220. Ziegler, C. G. P.; Hennersdorf, F.; Weigand, J. J.; Wolf, R. Z. Anorg. Allg. Chem. 2020, 646, 552–557. Hoidn, C. M.; Rödl, C.; McCrea-Hendrick, M. L.; Block, T.; Pöttgen, R.; Ehlers, A. W.; Power, P. P.; Wolf, R. J. Am. Chem. Soc. 2018, 140, 13195–13199. Coburger, P.; Demeshko, S.; Rödl, C.; Hey-Hawkins, E.; Wolf, R. Angew. Chem. Int. Ed. 2017, 56, 15871–15875. Wolf, R.; Ehlers, A. W.; Khusniyarov, M. M.; Hartl, F.; de Bruin, B.; Long, G. J.; Grandjean, F.; Schappacher, F. M.; Pöttgen, R.; Slootweg, J. C.; Lutz, M.; Spek, A. L.; Lammertsma, K. Chem. Eur. J. 2010, 16, 14322–14334. Wolf, R.; Ghavtadze, N.; Weber, K.; Schnöckelborg, E.-M.; de Bruin, B.; Ehlers, A. W.; Lammertsma, K. Dalton Trans. 2010, 39, 1453–1456. Wolf, R.; Slootweg, J. C.; Ehlers, A. W.; Hartl, F.; de Bruin, B.; Lutz, M.; Spek, A. L.; Lammertsma, K. Angew. Chem. Int. Ed. 2009, 48, 3104–3107. Wolf, R.; Ehlers, A. W.; Slootweg, J. C.; Lutz, M.; Gudat, D.; Hunger, M.; Spek, A. L.; Lammertsma, K. Angew. Chem. Int. Ed. 2008, 47, 4584–4587. Maier, T. M.; Sandl, S.; Shenderovich, I. G.; Jacobi von Wangelin, A.; Weigand, J. J.; Wolf, R. Chem. Eur. J. 2019, 25, 238–245. Pelties, S.; Maier, T.; Herrmann, D.; de Bruin, B.; Rebreyend, C.; Gärtner, S.; Shenderovich, I. G.; Wolf, R. Chem. Eur. J. 2017, 23, 6094–6102. Büschelberger, P.; Gärtner, D.; Reyes-Rodriguez, E.; Kreyenschmidt, F.; Koszinowski, K.; Jacobi von Wangelin, A.; Wolf, R. Chem. Eur. J. 2017, 23, 3139–3151. Gärtner, D.; Welther, A.; Rad, B. R.; Wolf, R.; Jacobi von Wangelin, A. Angew. Chem. Int. Ed. 2014, 53, 3722–3726. Weber, K.; Schnöckelborg, E.-M.; Wolf, R. ChemCatChem 2011, 3, 1572–1577. Roy, S.; Mondal, K. C.; Roesky, H. W. Acc. Chem. Res. 2016, 49, 357–369. Romain, C.; Bellemin-Laponnaz, S.; Dagorne, S. Coord. Chem. Rev. 2020, 422, 213411. Li, Y.; Mondal, K. C.; Roesky, H. W.; Zhu, H.; Stollberg, P.; Herbst-Irmer, R.; Stalke, D.; Andrada, D. M. J. Am. Chem. Soc. 2013, 135, 12422–12428. Arrowsmith, M.; Braunschweig, H.; Celik, M. A.; Dellermann, T.; Dewhurst, R. D.; Ewing, W. C.; Hammond, K.; Kramer, T.; Krummenacher, I.; Mies, J.; Radacki, K.; Schuster, J. K. Nature Chem. 2016, 8, 638–642. Singh, A. P.; Samuel, P. P.; Roesky, H. W.; Schwarzer, M. C.; Frenking, G.; Sidhu, N. S.; Dittrich, B. J. Am. Chem. Soc. 2013, 135, 7324–7329. Weinberger, D. S.; Melaimi, M.; Moore, C. E.; Rheingold, A. L.; Frenking, G.; Jerabek, P.; Bertrand, G. Angew. Chem. Int. Ed. 2013, 52, 8964–8967. Gstöttmayr, C. W. K.; Böhm, V. P. W.; Herdtweck, E.; Grosche, M.; Herrmann, W. A. Angew. Chem. Int. Ed. 2002, 41, 1363–1365. Arduengo, A. J.; Gamper, S. F.; Calabrese, J. C.; Davidson, F. J. Am. Chem. Soc. 1994, 116, 4391–4394. Schaub, T.; Backes, M.; Radius, U. Organometallics 2006, 25, 4196–4206.

107

108

Very Low Oxidation States in Organometallic Chemistry

262. Lee, C. H.; Laitar, D. S.; Mueller, P.; Sadighi, J. P. J. Am. Chem. Soc. 2007, 129, 13802–13803. 263. Mondal, K. C.; Samuel, P. P.; Li, Y.; Roesky, H. W.; Roy, S.; Ackermann, L.; Sidhu, N. S.; Sheldrick, G. M.; Carl, E.; Demeshko, S.; De, S.; Parameswaran, P.; Ungur, L.; Chibotaru, L. F.; Andrada, D. M. Eur. J. Inorg. Chem. 2014, 2014, 818–823. 264. Weinberger, D. S.; Amin Sk, N.; Mondal, K. C.; Melaimi, M.; Bertrand, G.; Stückl, A. C.; Roesky, H. W.; Dittrich, B.; Demeshko, S.; Schwederski, B.; Kaim, W.; Jerabek, P.; Frenking, G. J. Am. Chem. Soc. 2014, 136, 6235–6238. 265. Du, J.; Chen, W.; Chen, Q.; Leng, X.; Meng, Y.-S.; Gao, S.; Deng, L. Organometallics 2020, 39, 729–739. 266. Wang, D.; Chen, Q.; Leng, X.; Deng, L. Inorg. Chem. 2018, 57, 15600–15609. 267. Ung, G.; Rittle, J.; Soleilhavoup, M.; Bertrand, G.; Peters, J. C. Angew. Chem. Int. Ed. 2014, 53, 8427–8431. 268. Paul, U. S. D.; Radius, U. Organometallics 2017, 36, 1398–1407. 269. Kretschmer, R.; Ruiz, D. A.; Moore, C. E.; Rheingold, A. L.; Bertrand, G. Angew. Chem. Int. Ed. 2014, 53, 8176–8179. 270. Jones, C.; Sidiropoulos, A.; Holzmann, N.; Frenking, G.; Stasch, A. Chem. Commun. 2012, 48, 9855–9857. 271. Melancon, K. M.; Gildner, M. B.; Hudnall, T. W. Chem. Eur. J. 2018, 24, 9264–9268. 272. Abraham, M. Y.; Wang, Y.; Xie, Y.; Wei, P.; Schaefer, H. F.; Schleyer, P. V. R.; Robinson, G. H. Chem. Eur. J. 2010, 16, 432–435. 273. Mondal, K. C.; Samuel, P. P.; Roesky, H. W.; Carl, E.; Herbst-Irmer, R.; Stalke, D.; Schwederski, B.; Kaim, W.; Ungur, L.; Chibotaru, L. F.; Hermann, M.; Frenking, G. J. Am. Chem. Soc. 2014, 136, 1770–1773. 274. Gao, Y.; Li, G.; Deng, L. J. Am. Chem. Soc. 2018, 140, 2239–2250. 275. Ung, G.; Peters, J. C. Angew. Chem. Int. Ed. 2015, 54, 532–535. 276. Rudd, P. A.; Planas, N.; Bill, E.; Gagliardi, L.; Lu, C. C. Eur. J. Inorg. Chem. 2013, 2013, 3898–3906. 277. Hammer, R.; Klein, H.-F.; Friedrich, P.; Huttner, G. Angew. Chem. Int. Ed. 1977, 16, 485–486. 278. Klein, H. F.; Koenig, H.; Koppert, S.; Ellrich, K.; Riede, J. Organometallics 1987, 6, 1341–1345. 279. Yamamoto, A.; Miura, Y.; Ito, T.; Chen, H. L.; Iri, K.; Ozawa, F.; Miki, K.; Sei, T.; Tanaka, N.; Kasai, N. Organometallics 1983, 2, 1429–1436. 280. Anderson, D. M.; Cloke, F. G. N.; Cox, P. A.; Edelstein, N.; Green, J. C.; Pang, T.; Sameh, A. A.; Shalimoff, G. J. Chem. Soc., Chem. Commun. 1989, 53–55. 281. Cloke, F. G. N. Chem. Soc. Rev. 1993, 22, 17. 282. Brennan, J. G.; Cloke, F. G. N.; Sameh, A. A.; Zalkin, A. J. Chem. Soc., Chem. Commun. 1987, 1668–1669. 283. Bochkarev, M. N.; Trifonov, A. A.; Fedorova, E. A.; Emelyanova, N. S.; Basalgina, T. A.; Kalinina, G. S.; Razuvaev, G. A. J. Organomet. Chem. 1989, 372, 217–224. 284. Selikhov, A. N.; Cherkasov, A. V.; Fukin, G. K.; Trifonov, A. A.; del Rosal, I.; Maron, L. Organometallics 2015, 34, 555–562. 285. Edelmann, A.; Blaurock, S.; Lorenz, V.; Hilfert, L.; Edelmann, F. T. Angew. Chem. 2007, 119, 6855–6857. 286. Huang, W.; Diaconescu, P. L. Dalton Trans. 2015, 44, 15360–15371. 287. Kelly, R. P.; Maron, L.; Scopelliti, R.; Mazzanti, M. Angew. Chem. Int. Ed. 2017, 56, 15663–15666. 288. Bochkarev, M. N. Chem. Rev. 2002, 102, 2089–2118. 289. Evans, W. J. Organometallics 2016, 35, 3088–3100. 290. Kelly, R. P.; Toniolo, D.; Tirani, F. F.; Maron, L.; Mazzanti, M. Chem. Commun. 2018, 54, 10268–10271. 291. Selikhov, A. N.; Mahrova, T. V.; Cherkasov, A. V.; Fukin, G. K.; Kirillov, E.; Alvarez Lamsfus, C.; Maron, L.; Trifonov, A. A. Organometallics 2016, 35, 2401–2409. 292. Huang, W.; Dulong, F.; Wu, T.; Khan, S. I.; Miller, J. T.; Cantat, T.; Diaconescu, P. L. Nat. Commun. 2013, 4, 1448. 293. Langeslay, R. R.; Fieser, M. E.; Ziller, J. W.; Furche, F.; Evans, W. J. Chem. Sci. 2015, 6, 517–521. 294. Langeslay, R. R.; Fieser, M. E.; Ziller, J. W.; Furche, F.; Evans, W. J. J. Am. Chem. Soc. 2016, 138, 4036–4045. 295. MacDonald, M. R.; Fieser, M. E.; Bates, J. E.; Ziller, J. W.; Furche, F.; Evans, W. J. J. Am. Chem. Soc. 2013, 135, 13310–13313. 296. Billow, B. S.; Livesay, B. N.; Mokhtarzadeh, C. C.; McCracken, J.; Shores, M. P.; Boncella, J. M.; Odom, A. L. J. Am. Chem. Soc. 2018, 140, 17369–17373. 297. Windorff, C. J.; Chen, G. P.; Cross, J. N.; Evans, W. J.; Furche, F.; Gaunt, A. J.; Janicke, M. T.; Kozimor, S. A.; Scott, B. L. J. Am. Chem. Soc. 2017, 139, 3970–3973. 298. Su, J.; Windorff, C. J.; Batista, E. R.; Evans, W. J.; Gaunt, A. J.; Janicke, M. T.; Kozimor, S. A.; Scott, B. L.; Woen, D. H.; Yang, P. J. Am. Chem. Soc. 2018, 140, 7425–7428. 299. Niu, S.; Cai, H.-X.; Zhao, H.-B.; Li, L.; Pan, Q.-J. RSC Adv. 2020, 10, 26880–26887. 300. Tian, J.-N.; Zheng, M.; Li, L.; Schreckenbach, G.; Guo, Y.-R.; Pan, Q.-J. New J. Chem. 2019, 43, 1469–1477. 301. Fang, B.; Ren, W.; Hou, G.; Zi, G.; Fang, D.-C.; Maron, L.; Walter, M. D. J. Am. Chem. Soc. 2014, 136, 17249–17261. 302. Diaconescu, P. L.; Arnold, P. L.; Baker, T. A.; Mindiola, D. J.; Cummins, C. C. J. Am. Chem. Soc. 2000, 122, 6108–6109. 303. Mills, D. P.; Moro, F.; McMaster, J.; van Slageren, J.; Lewis, W.; Blake, A. J.; Liddle, S. T. Nat. Chem. 2011, 3, 454–460. 304. Monreal, M. J.; Khan, S. I.; Kiplinger, J. L.; Diaconescu, P. L. Chem. Commun. 2011, 47, 9119–9121. 305. Evans, W. J.; Traina, C. A.; Ziller, J. W. J. Am. Chem. Soc. 2009, 131, 17473–17481. 306. Evans, W. J.; Kozimor, S. A.; Ziller, J. W.; Kaltsoyannis, N. J. Am. Chem. Soc. 2004, 126, 14533–14547. 307. Zhang, L.; Fang, B.; Hou, G.; Ai, L.; Ding, W.; Walter, M. D.; Zi, G. Dalton Trans. 2016, 45, 16441–16452. 308. Zhang, L.; Hou, G.; Zi, G.; Ding, W.; Walter, M. D. J. Am. Chem. Soc. 2016, 138, 5130–5142. 309. Taylor, J. W.; McSkimming, A.; Moret, M.-E.; Harman, W. H. Angew. Chem. Int. Ed. 2017, 56, 10413–10417. 310. Raubenheimer, H. G.; Schmidbaur, H. Organometallics 2012, 31, 2507–2522. 311. Hicks, J.; Mansikkamäki, A.; Vasko, P.; Goicoechea, J. M.; Aldridge, S. Nat. Chem. 2019, 11, 237–241.

1.05

Very High Oxidation States in Organometallic Chemistry

Moritz Malischewski, Freie Universität Berlin, Institut für Chemie und Biochemie—Anorganische Chemie, Berlin, Germany © 2022 Elsevier Ltd. All rights reserved.

1.05.1 Introduction 1.05.2 Metal alkyl complexes 1.05.3 Metal aryl complexes 1.05.4 Alkyl- and aryl complexes with oxo, imido and nitrido ligands 1.05.5 Carbenes and carbynes 1.05.6 Cyclopentadienyl complexes 1.05.7 Hydride 1.05.8 Silyl chemistry 1.05.8.1 Halogenation reactions 1.05.9 One-electron oxidizing agents 1.05.10 Conclusion Acknowledgement References

109 110 113 118 121 124 125 127 128 129 131 131 131

Abbreviations AgRE Ar BArF24 Biphe Cp Cp dbabh Dipp Dippe dmp Et IMes L M mCPBA Me Mes NBu4 NHC NHE Nor NTf2 OAcF OTf Ph PPh4 R SIMes Solv tBu THT Ts

Silver reference electrode Aryl Tetrakis(3,5-bis(trifluoromethyl)phenyl)borate Biphenyl-2,20 -diyl Cyclopentadienyl Pentamethylcyclopentadienyl 2,3:5,6-Dibenzo-7-aza bicyclo[2.2.1]hepta-2,5-diene 2,6-Diisopropylphenyl 1,2-Bis-(diisopropylphosphino)ethane N,N0 -dimethylpiperazine Ethyl 1,3-Dimesitylimidazol-2-ylidene Ligand Metal Meta-chloroperbenzoic acid Methyl Mesityl Tetrabutylammonium N-heterocyclic carbene Normal hydrogen electrode Norbornyl Triflimide Trifluoroacetate Triflate Phenyl Tetraphenylphosphonium Organic residue 1,3-Bis(2,4,6-trimethylphenyl)-4,5-dihydroimidazol-2-ylidene Solvent Tertbutyl Tetrahydrothiophene Tosylate

Comprehensive Organometallic Chemistry IV

https://doi.org/10.1016/B978-0-12-820206-7.00004-4

109

110

1.05.1

Very High Oxidation States in Organometallic Chemistry

Introduction

The formal oxidation state of a metal atom in a complex is the hypothetical charge if the bonding electron pairs were assigned to the more electronegative binding partner (all its bonds were 100% ionic). While this concept is a useful formalism for bookkeeping and to track electron transfer in redox reactions, it should be taken with a grain of salt. For example, high formal oxidation states are not necessarily associated with strong oxidation power. For example, ReF7, ReO−4 and ReH2− 9 are all formally rhenium(VII) compounds. Of these, only ReF7 is a very powerful oxidant while ReO−4 is mildly oxidizing, at best. ReH2− 9 is an exceptional example, because its rhenium(VII) oxidation state belies the fact that it is formed by reduction of ReO−4 with alkali metals in ethanol.1 The preparation of organometallic compounds in very high oxidation states is an important research area in organometallic chemistry. However, the related experimental work is often challenging since two main problems have to be addressed: 1. The target molecule could be intrinsically unstable, since various decomposition pathways are available to organometallic species (e.g. reductive elimination, b-hydride elimination, H atom abstraction from the solvent, homolytic metal-carbon bond cleavage with formation of radicals).2 2. Is there a viable synthetic route to prepare the target molecule in the desired oxidation state? This could involve the reaction of an organometallic compound in a relatively low oxidation state with an oxidant or the substitution of a labile ligand (e.g. halide) in a high valent metal complex by an organometallic nucleophile (e.g.: CH3Li, C6H5MgBr). In the latter, the oxidation state of the metal could remain unchanged, decrease (in case of strongly reducing organometallic reagents) or even increase in case of disproportionation reactions! Indeed, many examples of unusually high oxidation states covered in this review were prepared under non-oxidizing or even reducing conditions, which demonstrates that the proper choice of the transmetallating agent is of great importance. Since the term “very high oxidation state” is rather subjective, this review is not intended to be an encyclopedic list of compounds. Instead, it discusses strategies how organic ligands (alkyl, aryl) can be designed to be less susceptible to decomposition reactions and how inorganic co-ligands can affect the stability of high oxidation states. Additionally, different synthetic routes towards such high valent organometallic complexes are compared (oxidation vs transmetallation). A brief overview of powerful one-electron oxidizers is given at the end for practical advice in this regard. In general, the stability of high oxidation states increases in the series 3d < 4d < 5d, and higher oxidation states can be realized for 4d and 5d metals than for the first series of transition metals. Consequently, the chemistry of 3d metal complexes can differ significantly from their heavier homologues. Furthermore, the ability to access a given oxidation state depends significantly on the choice of the ligand system. For electrostatic reasons, high oxidation states are typically stabilized by anionic donor ligands, (alkyl R−, aryl R−, cyclopentadienyl Cp−, hydride H−, silyl R3Si−, halide X−, oxide O2−, imide NR2−, nitride N3−). The last three are powerful p-donors and are therefore optimal ligands for high-valent metal ions. Additionally, this review will also briefly discuss the chemistry of carbene (R2C]M) and carbyne (RC^M) ligands. One word of caution: oxidation states are a formalism. However, such a formal oxidation state may not be the correct description of the actual distribution of the bonding electrons, as revealed by spectroscopical measurements (e.g. non-innocent ligands). However, although the description of organometallic complexes containing strongly covalent bonds (MdH, MdC, MdSi) with the oxidation state formalism (which is based on purely ionic attribution of electrons) leads to inaccuracies, formal oxidation states remain a helpful concept for the synthetic chemist who thinks in cationic or anionic synthons or tries to categorize a new compound.

1.05.2

Metal alkyl complexes

The stability of metal-alkyl complexes depends on such factors as oxidation state, coordination number, d-electron count and steric protection. Isolating these complexes requires minimizing the feasibility and rates of decomposition pathways, and several common ones are outlined here. Various decomposition pathways are possible with transition metal alkyls, including reductive elimination and b-hydride elimination as well as homolytic metal-carbon bond cleavage with formation of radicals.2 The latter is linked to metal-carbon bond strength. In general, 4d and 5d metals form stronger M-C bonds than 3d metals.3 Electronegative substituents can increase the metal-carbon bond strength; for example MdCF3 bonds are typically stronger than MdCH3 bonds and therefore less susceptible to homolysis.4 The rate of reductive elimination of alkyl groups to form carbon-carbon bonds depends on the energy of the transition state, in which the alkyl groups must bend. This canting of the ligands can be prevented using steric hindrance between two bulky alkyl groups. Additionally, steric saturation prevents the existence of free coordination sites, which might trigger decomposition reactions. b-Hydride elimination as a decomposition pathway is typically avoided by choosing ligands without b-hydrogens. As a consequence bulky ligands have been developed where b- (and also a-) hydride abstraction is impossible/disfavored (Scheme 1).

Scheme 1 Bulky alkyl ligands where b-hydride elimination is unfavorable.

Very High Oxidation States in Organometallic Chemistry

111

In 1972 Bower and Tennent prepared a series of tetrakis(1-norbornyl) complexes M(Nor)4 (M ¼ Ti, Zr, Hf, V, Cr, Mn, Fe, Co).5 While the tetravalent chromium complex is by far the most stable to thermolysis, diluted acids or air, all compounds could be isolated. The general synthetic route involves the reaction of 1-norbornyllithium with metal halides (in various oxidation states: MnBr2, CoCl2, FeCl3, CrCl3, TiCl4, ZrCl4, HfCl4, VCl4) in pentane with tumbling glass beads. The di- and trivalent metal ions probably form anionic species which undergo disproportionation or oxidation to give the tetravalent complexes in the end. Due to the bulkiness and strong-s-donor properties of the norbornyl ligands and the inaccessibility of b-hydride elimination pathways, unusually high oxidation states can be stabilized. Surprisingly, tetrakis(1-norbornyl)cobalt Co(Nor)4 can even be oxidized at relatively low potentials (−0.65 V vs Cp2Fe+/Cp2Fe, or chemically by Ag+) to Co(Nor)+4 containing formally pentavalent cobalt.6 For the tetravalent iron derivative Fe(Nor)4, the Mössbauer spectrum has been reported.7 Tetravalent nickel compounds (Nor)3NiBr are also known.8 Recently, a homoleptic Ni(IV) tetraalkyl complex (however with a different ligand) was reported (Scheme 2).9 In general, the unexpectedly high stability of the known high-valent norbornyl complexes is probably caused by a combination of steric and favorable dispersive interactions.10 London dispersion forces can actually be attractive in sterically crowded molecules at the appropriate H. . .H distances, which is an enthalpic stabilization. As a consequence, energy barriers (for example for decomposition reactions like reductive elimination) can differ significantly from less crowded molecules. The existence of FeR4 (R ¼ 2adamantyl, cyclohexyl) can be justified similarly.11

Scheme 2 Structure of a nickelaspirocyclononane, a tetraalkyl complex of nickel(IV).

The chemistry of the 4d elements with the 1-norbornyl ligand has been less explored. Mo(Nor)4 can be prepared in 25% yield from MoCl3(THF)3 and norbornyllithium, and has a triplet ground state (solution moment of 2.62 mB at 300 K).12 Cyclovoltammetric measurements in CH2Cl2 reveal two oxidation waves (−0.15 V and + 1.25 V vs Ag/Ag+), but only the first is reversible. Because chemical oxidation attempts with ferrocenium or Ag+ were reported to be unsuccessful, the putative molybdenum(V) alkyl remains unobserved. The major disadvantage of the 1-norbornyl group is its tendency towards disorder in crystal structures. Probably, this problem could be circumvented by using the more symmetric bicyclo[2.2.2]octyl group instead. Unfortunately, 1-bromobicyclo[2.2.2] octane is very expensive (>$1000 per gram). In terms of suppressing b-hydride eliminations and steric bulk, the 1-adamantyl-group should be even superior. Moreover, the corresponding halides are relatively cheap and commercially available in large scale. However, transition metal adamantyl chemistry is poorly developed due to problematic side reactions and very low solubility of homoleptic adamantyl-complexes in organic solvents.13 The 1-adamantylmethyl ligand is another possible alternative that has been underexplored.14 In addition, Pt, Pd and Re complexes of a perfluorinated norbornyl ligand are known, suggesting that other fluorinated analogues could be future targets.15 While the norbornyl group is the ligand of choice to stabilize very high oxidation states of 3d metals, 4d and 5d organometallic complexes in even higher oxidation states can be accessed with the simplest alkyl ligand: CH3! Hexamethyltungsten W(CH3)6 was first prepared by Wilkinson in 1973 by reaction of WCl6 with methyllithium.16 Later, an improved synthetic protocol was published which used trimethylaluminium instead (Scheme 3).17 Although W(CH3)6 decomposes slowly at room temperatures, several unexpected explosions of this compound were reported. Initially, the structure of W(CH3)6 was supposed to be octahedral. In 1996, Seppelt succeeded with the crystal structure determination of W(CH3)6, revealing a strongly distorted trigonal prismatic structure.18 This unusual structure can be rationalized by assuming a sd5 hybridization with bond angles of  63 and 117 since the sixcoordinated d0 metal center does not contain any p-donor ligands (which could otherwise donate into the empty t2g orbitals in an octahedral complex).19 The synthesis of Mo(CH3)6 is especially difficult, due to the higher oxidation power of Mo(VI). Consequently, the reaction between the powerful oxidant MoF6 and methyllithium is conducted in diethyl ether first at −130  C, and later at −78  C for 12 h.20 For the preparation of Re(CH3)6 a different synthetic procedure is used: ReO(CH3)4 is reacted with trimethylaluminium.17b,21 Interestingly, [M(CH3)6] (M ¼ Mo, W, Re) react with excess methyllithium to form anionic [Mo(CH3)7]−.20 [W(CH3)7]− 22 and [Re(CH3)8]2− (Scheme 3).22 The synthesis of [W(CH3)8]2− was claimed with incomplete evidence,17a and has been difficult to reproduce.22

Scheme 3 Preparation of homoleptic high-valent metal alkyls.

112

Very High Oxidation States in Organometallic Chemistry

For late transition metals the stability of high-oxidation-state metal alkyls is limited due to the ease of reductive elimination. This can be attributed to their higher electronegativities, and the correspondingly larger driving force for reduction. For example, trimethylgold Au(CH3)3 (perhaps its ether adduct) can be prepared in diethyl ether solution at −65  C, but decomposes above −40  C to elemental gold, ethane and methane (Scheme 4).23

Scheme 4 Preparation and decomposition of trimethylgold.

Moreover, the tetramethylaurate [Au(CH3)4]− is known.24 Attempts to oxidize this complex with one-electron oxidizers like ferrocenium, arenediazonium and nitrosonium cations lead to the formation of methane, ethane and elemental gold. In the presence of ligands like triphenylphosphine, the putative organogold(IV) species Au(CH3)4 may be formed but then rapidly decomposes via an homolytic pathway to CH3 radical and Au(CH3)3 which can be trapped as a complex with triphenylphosphine [PPh3Au(CH3)3].24a With group 11 metals, the trifluoromethyl group CF3 is a far better ligand for stabilizing metals in high oxidation states, forming stable complexes [M(CF3)4]− with M ¼ Cu, Ag and Au.25 Though it is counterintuitive that electron-withdrawing groups would stabilize high-valent complexes, this may arise mostly from an influence on the reaction kinetics. Recent spectroscopic studies have disputed the description of the copper in [Cu(CF3)4]− as +3, because this is a system in which the relative metal and ligand orbital energies are inverted relative to most metal complexes (see below).26 This highlights the difference between formal oxidation states (from conventions placing metal-ligand bonding electrons on the ligand) and “physical” oxidation states (from spectroscopic investigations judging the electron density on the metal). In fact, for most complexes that are formally Cu(III), their “physical” oxidation state is not Cu(III).27 In addition to the four-coordinate anions, neutral Au(CF3)3 has been mentioned in the literature. It was synthesized by co-condensation of trifluoromethyl radicals and gold vapor in a preparative scale and characterized in solution via 19F NMR spectroscopy. Decomposition to gold fluorides was observed at 0  C.28 Due to the lower stability of the oxidation state +3 for silver in comparison with gold,29 the existence of an organometallic Ag(III) complex in form of [Ag(CF3)4]− is especially surprising. However, its chemistry remained rather unexplored for several years. This might be caused by its original synthesis from the highly toxic Cd(CF3)2 and Ag+ salts.25c Recently, a significantly improved procedure using commercially available Me3SiCF3 was reported.30 Here oxidation of Ag(I) to Ag(III) was accomplished with a hypervalent iodine reagent (Scheme 5).

Scheme 5 Preparation of high-valent organosilver compounds, OAcF ¼ trifluoroacetate.

Although reductive elimination of CF3-CF3 from [M(CF3)4]− (M ¼ Cu, Ag, Au) is significantly exergonic, activation barriers for such a concerted process are very high. Moreover, heterolytic cleavage of M-CF3 is also energetically unfavorable. Instead, the preferred decomposition pathways of these compounds is via homolytic cleavage of MdCF3 bonds with formation of CF3 radicals.26 Consequently, the higher bond energy of MdCF3 bonds instead of MdCH3 groups is beneficial for the stability of such compounds.3 The reverse of reductive elimination, oxidative addition, is well known to give access to organometallic complexes if alkyl halides are reacted with electron-rich metal centers.31 One example for organogold chemistry is that oxidative addition of CH3I to [Au(CF3)2]− produces trans-bis(trifluoromethyl)methyliodoaurate [(CF3)2Au(CH3)I]− (Scheme 6).32

Scheme 6 Oxidative addition of alkyl halides to give Au(III) compounds.

Interestingly, trifluoromethyl groups also can be introduced via oxidative addition reactions. Addition of gaseous trifluoromethyliodide CF3I to Au(I) complexes can give access to gold(III) trifluoromethyl complexes (Scheme 7).33 The reaction may proceed via a radical pathway, since the reaction could be quenched by addition of the radical scavenger galvinoxyl.

Very High Oxidation States in Organometallic Chemistry

113

Scheme 7 Oxidative addition of CF3I to give Au(III) compounds.

Another elegant way to incorporate CF3 groups into heteroleptic metal complexes is the use of the Yagupolskii34 or Umemoto reagents.35 These commercially available36 S-(trifluoromethyl)-diarylsulfonium salts act as electrophilic CF3 transfer reagents. Although these reagents are typically used for the trifluoromethylation of organic substrates, they have also been used with success for the preparation of metal complexes that contain CF3 groups. For example, Sanford showed that the Umemoto reagent can oxidize Ni(II) to a stable Ni(IV) trifluoromethyl complex (Scheme 8).37

Scheme 8 Preparation of a Ni(IV) trifluoromethyl complex.

1.05.3

Metal aryl complexes

In general, the metal-aryl bond is stronger than the metal-alkyl bond.38 Additionally, b-hydride eliminations are far less common, since this would form a benzyne. Various ways exist to introduce aryl ligands into metal complexes. Formally, three different synthetic approaches are precedented. First, oxidative addition (for example of an aryl halide) or formal transfer of an aryl cation to the metal. Secondly, reaction between a metal complex and aryl radicals. Finally, substitution reactions involving formal aryl anions. Oxidative addition of carbon-halogen bonds to Pt(II) is well-known, however also carbon-hydrogen bonds can be activated to give Pt(IV) complexes (Scheme 9). Interestingly, when pentafluorophenyl groups are utilized, the corresponding arylplatinum(IV) fluorides can be isolated, demonstrating that even CdF bonds can oxidatively add to a metal center.39 Due to the lower stability of Pd(IV), isolation of such organometallics can be difficult.40 However, intramolecular oxidative addition of an aryl iodide moiety to a Pd(II) complex, including the crystallographic characterization of the corresponding unstable Pd(IV) aryl complex, was reported.41

Scheme 9 Oxidative addition of carbon-halogen and carbon-hydrogen bonds to Pt(II).

114

Very High Oxidation States in Organometallic Chemistry

Similarly, the isolation of organometallic Ni(IV) compounds is challenging. Phenyldiazonium ions, as well as diaryliodonium salts (as highly reactive formal sources of aryl cations) can be used for the electrophilic arylation of Ni(II) complexes to give Ni(IV) complexes (Scheme 10).42

Scheme 10 Electrophilic arylation of Ni(II) to give Ni(IV).

Moreover, aryldiazonium salts can be used for the formal transfer or aryl radicals when combined with a reducing agent like ferrocene. Sanford showed that Ni(III) complexes can be oxidatively arylated by carbon-centered radicals to give Ni(IV) complexes (Scheme 11).43 Alternatively, thermolysis of aryl peroxides can also be used to generate aryl radicals.

Scheme 11 Reaction of Ni(III) with aryl radical sources to produce Ni(IV).

While the aforementioned methods are suitable to introduce one aryl group into a metal complex, they cannot be used to prepare homoleptic metal aryl complexes. Homoleptic metal aryl complexes are typically prepared by reaction of high-valent metal halides with transmetallating agents (e.g. organolithium, organomagnesium, etc.) as with the corresponding metal alkyl complexes. These reactions may lead to complex mixtures due to side reactions, especially reductions of the metal center. Consequently, the choice of the transmetallating agent (and its reducing properties) are of great importance for the success of a reaction. Similar to the corresponding alkyl compounds, aryllithium reagents are more reactive than arylmagnesium reagents, but also stronger reductants, but other generalities are difficult to identify. For preparing homoleptic metal aryl complexes, unsubstituted phenyl groups work poorly:44 for example, tetraphenyltitanium Ti(C6H5)4 decomposes at −10  C to Ti(C6H5)2 and biphenyl.45 In contrast, the far less air-sensitive tetramesityltitanium Ti(Mes)4 (mesityl ¼ 2,4,6-trimethylphenyl) has a decomposition temperature of +160  C!46 Interestingly, the pentafluorophenyl derivative Ti(C6F5)4 decomposes above +100  C.47 The ability of both electron-donating and electron-withdrawing substituents to stabilize these species suggests that steric hindrance to accessing the reductive elimination transition state is the most important factor. In addition, the ortho CdH are susceptible to loss resulting in benzyne formation, and so incorporating ortho substituents is beneficial. The greater bond strength of MdC6F5 in contrast to MdC6H5 is also beneficial to suppress decomposition via homolysis of the MdC bonds and formation of radicals. As already mentioned, the mesityl (2,4,6-Me3C6H2) group is a privileged ligand in organometallic chemistry. It was used extensively by Wilkinson to prepare a series of spectacular high-valent organometallic compounds M(Mes)3 (M ¼ Rh, Ir) and M(Mes)4 (M ¼ Ru, Ir).48 Trigonal-pyramidal Rh(Mes)3 and Ir(Mes)3 come from the reaction of tetrahydrothiophene adducts of the corresponding metal trichlorides MCl3(tht)3 with mesityl Grignard reagents (Scheme 12).49 The reaction of IrCl3 ∙ nH2O with mesityllithium gives Ir(Mes)4, albeit in only 15% yield.50 However, the degree of hydration of the starting material is of utmost importance.48,49b Ru(Mes)4 is similarly produced from RuCl3(tht)3 and (Mes)2Mg(THF)2 in 18% yield.49b The low yields are worrisome, and even Wilkinson stated: “The success (or failure) in getting the expected products in this type of metal chemistry is a matter of trial and is unpredictable: the reasons and mechanisms are also far from clear. [. . .] In short one can say that this chemistry hardly qualifies as science – more of an art form in fact.”48

Very High Oxidation States in Organometallic Chemistry

115

Scheme 12 Preparation of homoleptic mesityliridium compounds

Unexpectedly, these mesityl complexes can be oxidized further, as cyclic voltammetry of trimesitylrhodium Rh(Mes)3 displays two oxidation waves: the first is reversible at +0.75 V (vs Cp2Fe+/Cp2Fe) and a second is irreversible at slightly higher potential, however chemical oxidation attempts with AgBF4 were unsuccessful (probably due to the insufficient oxidation power of Ag+).51 In contrast, Ir(Mes)3 is extremely air-sensitive and reacts with dioxygen with formation of the Ir(V)-oxo complex Mes3IrO (Scheme 13).52 Interestingly, the reaction of Ir(Mes)3 with mesityl azide MesN3 did not give the expected Iridium(V)-imido complex (Mes)3IrNMes, but an Ir(III)-tetraazenido complex was obtained in 67% yield.53 The iridium center is coordinated to three nitrogen atoms as well as an olefin in trans-configuration, which is probably formed upon dehydrogenative coupling of two mesityl groups.

Scheme 13 Oxidation reactions of Ir(Mes)3.

Oxidation of stable Ir(Mes)4 and Ru(Mes)4 to the corresponding cations occurs at potentials of −0.44 V and + 0.32 V (vs Cp2Fe+/ Cp2Fe).49b While the former can be oxidized by Ag+ or NO+, oxidation of the ruthenium compound is only possible with NO+. Triflate or hexafluorophosphate salts of the pentavalent ruthenium and iridium compounds are highly stable and have decomposition temperatures >200  C. [Ir(Mes)4]+ CF3SO−3
CH3CN has been structurally characterized while recently the crystal structure of a similar pentavalent ruthenium compound was determined ([RuR4]+ BF−4; R ¼ 4-methoxy-2-methylphenyl).54 Tetrakis(tolyl)osmium Os(2-MeC6H4)4 in CH2Cl2 displays one reversible oxidation wave at +0.33 V (vs Cp2Fe+/Cp2Fe). Accordingly, reaction with AgOTf gave crystalline [Os(2-MeC6H4)4]+ OTf−.55 In general, the redox potential of the ruthenium compounds seems to be only marginally higher than of the corresponding osmium compounds.56 In 2020, a better synthetic route to Os(2-MeC6H4) was reported.57 Instead of the original procedure that used OsO4 and 2-tolylmagnesium bromide,58 (NBu4)2[OsBr6] was reacted with arylmagnesium reagents. While the 2-tolyl and the 2,5-xylyl derivative were formed in  75% yield, the so far unknown 2,4,6-mesityl derivative could be isolated in 21% yield (for ruthenium this method worked less well). Cyclic voltammetry for Os(Mes)4 revealed two reversible one-electron oxidations in CH2Cl2 at +0.15 V and +1.12 V vs [Cp2Fe+]/ [Cp2Fe]. The comparison of the electrochemical data for the tetraarylosmium compounds indicated that for each additional methyl group the oxidation potentials decrease by 22 mV. It is unknown whether the second oxidation of the ruthenium and iridium compounds would be at significantly higher electrochemical potentials than for the osmium compounds. Additionally, it is unclear if such dications would be stable enough to be isolated. The existence of a purely organometallic Ir(VI) compound would be truly remarkable. One could imagine that the oxidation state +6 could be even better stabilized if the mesityl groups (C6H2Me3) could be replaced by pentamethylphenyl (C6Me5) groups. In 2002 it was stated that the C6Me5 group is a “relevant, but possibly overlooked ligand in organometallic chemistry”, since its chemistry is virtually unexplored.59

116

Very High Oxidation States in Organometallic Chemistry

As described earlier, homoleptic trifluoromethyl complexes [M(CF3)4]− have been isolated for M ¼ Cu, Ag and Au. The corresponding pentafluorophenyl complexes [M(C6F5)4]− exist only for gold, since the oxidation state +3 is most stable for it.60 It is prepared by reacting the tetrahydrothiophene (tht) adduct of AuCl3 with pentafluorphenyllithium (Scheme 14), but this gives partial reduction. Unexpectedly, even the unfluorinated Au(III) species NBu+4 [Au(C6H5)4]− exists.61

Scheme 14 Preparation of homoleptic pentafluorophenyl complexes of Au(I) and Au(III).

The only examples of tetraarylcopper(III) complexes use two biphenyl-2,20 -diyl (Biphe) ligands. The crystal structure of [Li(THF)4]+ [Cu(Biphe)2]− reveals a distorted square-planar geometry for Cu(III), while the corresponding Cu(I) compound is tetrahedral. The synthesis involves the oxidation of a copper(I) species with benzoquinone (Scheme 15).62 So far, no tetraarylsilver(III) species is known.

Scheme 15 Preparation of a homoleptic Cu(III) aryl complex.

For group 10, complexes of the general formula [M(C6X5)4]2− have been studied (X ¼ F, Cl, M ¼ Ni, Pd, Pt). While for palladium only [Pd(C6F5)]2− is known,63 the corresponding Ni(III) complex NBu+4 [Ni(C6Cl5)4]− has been structurally characterized. It is prepared by oxidation of the Ni(II) compound with elemental chlorine.64 Later, NBu+4 [Ni(C6F5)4]− was also prepared through oxidation at relatively low potentials about +0.07 V (for C6F5) and −0.11 V (for C6Cl5) vs Cp2Fe+/Cp2Fe.65 Further oxidation to Ni(IV) could not be observed in the solvent window of dichloromethane. In contrast, at very high potentials, even Pt(IV) becomes accessible (Scheme 16). Pt(C6Cl5)4 is the only known homoleptic tetraarylplatinum complex66 which interestingly can be reduced to the oxidation state +3 as a stable monoanion.67

Scheme 16 Preparation of a homoleptic Pt(IV) aryl complex.

Since PtR4 would formally be a four-coordinate 14-electron complex, the existence of such a compound would be very unusual. However, the C6Cl5 ligand also offers the advantage of providing additional weak donor atoms. Indeed, the platinum atom in Pt(C6Cl5)4 displays a distorted octahedral coordination geometry (Scheme 17).66 Two platinum-chlorine distances of PtdCl 2.559 (2) and 2.681 (2) A˚ in the crystal structure indicate this interaction. Due to the higher polarizability of Cl in comparison to F, C6Cl5 should be superior to C6F5 as a ligand, because F bound to C is a only a very poor nucleophile. Coordination of ortho-chlorine atoms has also been observed for 3d transition metal complexes.68

Cl

Cl

Cl

Cl

Cl

Cl

Cl

Cl

Cl

Cl

Pt Cl

Cl

Cl

Cl Cl

Cl

Cl

Cl Cl

Scheme 17 Molecular Structure of Pt(C6Cl5)4 in the crystal structure of Pt(C6Cl5)4
Toluene.

Cl

Very High Oxidation States in Organometallic Chemistry

117

In summary, both the mesityl and the pentachlorophenyl ligand are both well suited for the preparation of high-valent metal complexes. They are both relatively similar in size, but have opposite electronic properties. The mesityl group is more electron donating, while the pentachlorophenyl is electron withdrawing, but provides extra binding groups.69 A related tetraarylplatinum(IV) complex nicely illustrates another powerful strategy to suppress reductive elimination. Instead of four aryl groups, two bidentate biphenyl-2,20 -diyl ligands are bound to platinum, while two additional tetrahydrothiophene ligands are also coordinated to give a 18-electron complex.70 The dianionic biphenyl-2,20 -diyl ligands additionally contain CF3 groups to suppress reductive elimination (Scheme 18). Furthermore, reductive elimination would result in a strained fourmembered ring (biphenylene). Fully halogenated biphenyl-2,20 -diyl ligands have received only very little attention.71

F3C

CF3 CF 3 S

Pt

S

F3C Scheme 18 Molecular Structure of cis-bis(tetrahydrothiophene)-bis(4,40 -bis(trifluoromethyl)-biphenyl-2,20 -diyl)-platinum(iv).

In general, biphenyl-2,20 -diyl (Biphe) ligands are not very bulky. Zirconium complexes with three [Zr(Biphe)3]2− or four [Zr(Biphe)4]4− of these ligands have been prepared and structurally characterized,72 which suggests that the use of Biphe ligands may be a general strategy to stabilize high valent transition metal complexes. At this point, it should be mentioned that the use of the biphe ligand has also led to spectacular discoveries in main group organometallic chemistry. In group 16, the thermal stability of tetraaryl-derivatives dramatically decreases in the order R4Te > R4Se > R4S.73 While Ph4S decomposes even at very low temperatures via reductive elimination (to Ph2S and biphenyl), the stability of the corresponding perfluorinated compound (C6F5)4S is limited to temperatures below 0  C.74 In contrast, for (Biphe)2S decomposition occurs at its melting point at +114  C (Scheme 19).75 In addition, even stable derivatives with hexavalent sulfur have been reported, namely: S(Biphe)2(CH3)2, S(Biphe)2(Ph)276 and [S(Biphe)2]2+ [BF4]−2.77 These findings illustrate the highly stabilizing effect of Biphe ligands on otherwise unstable molecules in high oxidation states.

Scheme 19 Highly unstable Ph4S (A) vs stable (Biphe)2S (B).

Despite the existence of W(CH3)6 for almost 50 years, there have been no examples of hexavalent homoleptic metal aryl complexes MR6, even though metal-aryl bonds are tendentially stronger than metal-alkyl bonds. Since the X-ray structures of several anionic complexes of group 4,72,78 578b,79 and 679b,80 metals with at least 6 phenyl rings are known, the inability to isolate a neutral species should not be primarily a steric problem. However, reductive elimination from such a high-valent metal center could occur easily. Additionally, reduction of M(VI) by organometallic reagents is another serious problem. For example, the reaction of WCl6 with C6F5Li in diethylether gives only Li[W(C6F5)5] which liberates very small amounts of [W(C6F5)5] upon thermal decomposition in high vacuum.81

118

Very High Oxidation States in Organometallic Chemistry

A particularly useful reagent for the transfer of C6F5 ligands without undesired reduction is the organothallium(III) reagent Tl(C6F5)2Cl. For example, it can convert [C6F5AuBr]− to [(C6F5)3AuBr]− and generates thallium chloride as a side product.82 Thus, the resistance of thallium to redox chemistry is beneficial. Whether the above-mentioned biphenyl-2,20 -diyl ligands could give access to a homoleptic W(VI) aryl complex is speculation. In the future, modification of the ligands could be used (introduction of electron-withdrawing substituents, blocking of position ortho to the metal-carbon bonds to avoid benzyne formation etc.).

1.05.4

Alkyl- and aryl complexes with oxo, imido and nitrido ligands

The oxo ligand is a powerful p-donor, and therefore it is well suited to stabilize metal complexes in high oxidation states. Two main synthetic pathways towards high-valent target molecules are (a) the treatment of a high oxidation state metal oxo complex with a transmetallating reagent or (b) the reaction of a low valent alkylmetal complex with an oxidizing source of oxygen. For the first method, the chemistry of rhenium and osmium is an instructive example. Methyltrioxorhenium CH3ReO3 is a high valent organometallic compound of significant importance in coordination chemistry and catalysis.83 It cannot be synthesized by reaction of ReO3X (X ¼ Cl, ReO4, OSiMe3) with organolithium, -magnesium or aluminum compounds since they are too strongly reducing. Instead Re2O7 is reacted with Sn(CH3)4 to give CH3ReO3 in quantitative yield (Scheme 20).84 This illustrates that the proper choice of the transmetallating agent is crucial for the preparation of high valent organometallic compounds. The reaction of Re2O7 with Zn(CH3)2 gives CH3ReO3 if the reaction is conducted at −78  C. Warming to −30  C or above leads to the formation of organorhenium(VI) compounds (CH3)4Re2O4 and (CH3)6Re2O3.84 Similarly, aryltrioxorhenium(VII) compounds are accessible via the organozinc reagents.85 However, the reaction depends significantly on the used stochiometries: 1:1 ratios of Re2O7 and ZnR2 can lead to the formation of R2ReO2 as the main product.86

Scheme 20 Preparation of CH3ReO3.

In contrast to rhenium, where transmetallation of Re2O7 is possible in some cases without reduction, the reaction of OsO4 with organometallic reagents inevitably leads to reduced species. For example, OsO(CH2SiMe3)4 is formed upon reaction of OsO4 with Mg(CH2SiMe3)2 in pentane at −70  C.87 Similarly, OsO(CH3)4 is formed in low yield by reaction of dimethylzinc with OsO4.88 Tetraarylosmium(IV) compounds OsR4 can be accessed via the reaction of OsO4 with triarylaluminium58 or arylmagnesium reagents.89 The latter can also form organoosmium(VI) compounds OsO2R2.90 Consequently, no organometallic compound in the oxidation state +8 is known so far. It is unclear whether imido-based Os(VIII) compounds as [Os(NtBu)4]91 or mixed oxo/imido species92 would be more promising starting materials than OsO4. The chemistry of imido complexes will be discussed later in this section. By starting from a low-valent complex, treatment with a suitable oxygen-transfer reagent (oxidant) can lead to the formation of a high-valent oxo complex.93 Various reagents have been used in the literature, for example dimethyldioxirane for the preparation of an organometallic platinum(IV) oxo complex (formation of acetone as a side product—Scheme 21).94 Ozone (O3) was recently used for the preparation of a Fe(IV) aqua-oxo complex,95 however it has not really found its way into organometallic chemistry. Peracids, e.g. meta-chloroperbenzoic acid (mCPBA) have been successfully used to generate a Fe(V) oxo complex.96 mCPBA, organic peroxides and aqueous H2O2 have been demonstrated to oxidize dimethylplatinum(II) complexes to platinum(IV), simultaneously introducing hydroxo and carboxylato ligands.97

Scheme 21 Preparation of an organometallic Pt(IV) oxo complex.

Very High Oxidation States in Organometallic Chemistry

119

Treatment of a dicationic Fe(II) tetracarbene complex with a iodosylbenzene derivative (RdI]O) as oxygen-transfer reagent gave a dicationic Fe(IV) oxo tetracarbene complex which was isolated and structurally characterized (Scheme 22).98 The introduction of substituents on the aryl ring of the hypervalent iodine reagent allows to tune the reactivity and to improve the solubility, and the corresponding aryl iodide is formed as by-product of the oxidation reaction.

Scheme 22 Preparation of a Fe(IV) oxo tetracarbene complex.

Imido (RN2−) and nitrido (N3−) ligands are also powerful p-donors and therefore well suited to stabilize metal centers in very high oxidation states. Additionally, these ligands have a lower tendency to be reduced by organometallic reagents than the corresponding oxo complexes. One instructive example is the preparation of an organomanganese compound in oxidation state +7. Mn(NtBu)3C6F5 can be prepared by reaction of Mn(NtBu)3Cl with AgC6F5 in Et2O (Scheme 23), but it slowly decomposes at room temperature and could only be characterized via NMR spectroscopy and mass spectrometry. However, it was the only compound with a Mn(VII)dC bond that could be successfully be prepared in this study.99 This extreme case of a Mn(VII) organometallic compound at the border of stability is probably only possible by the combination of three factors. First, the pentafluorophenyl ligand is well suited to form stable complexes in high oxidation states because of the strong bonds mentioned above. Additionally, the tert-butyl-imido ligand is a powerful p-donor that effectively stabilizes Mn(VII). In contrast to Mn(NtBu)3Cl,100 the corresponding oxo complex MnO3Cl is extremely unstable and sensitive to moisture and probably too strongly oxidizing.101

Scheme 23 Preparation of an organometallic Mn(VII) complex.

The use of tert-butyl-imido groups also allows the preparation of stable arylchromium(VI) species even using Grignard reagents (Scheme 24).102 Interestingly, no analogous compounds can be obtained when phenyl and 2-tolyl ligands are used instead of 2,6xylyl or 2,4,6-mesityl ligands. This reiterates the need to block both ortho-positions of the phenyl ring by methyl groups to suppress side reactions.

Scheme 24 Preparation of an organometallic Cr(VI) complex.

Another elegant way to prepare high-valent metal imido complexes are oxidation reactions with organic azides RN3. In principle, such reactions can proceed by extrusion of one molecule of nitrogen and the imido moiety is transferred to the metal. The oxidation state of the metal increases by +2. For example, a Fe(IV) imido complex can be prepared with this method (Scheme 25).103 Similar examples are further discussed in the carbene section. Surprisingly, the success of generating a high valent imido complex may depend on the choice of the organic azide, and their reactivity can differ significantly.104 Furthermore, side reactions (for example migratory insertion between imide NR and alkyl R’ moieties to give an amide ligand NRR’) may occur.105

120

Very High Oxidation States in Organometallic Chemistry

Scheme 25 Preparation of a Fe(IV) imido complex.

A somewhat similar reaction for the oxidative transfer of RN groups to an organometallic compound is the use of Chloramine T (N-chlorotosylamide, sodium salt). This reagent is a cheap and stable powerful oxidant.106 A few examples from organo main group chemistry and inorganic coordination chemistry exist, which demonstrate the usefulness of this reagent. Its reactivity with organotransition metal complexes seems to be less studied. Trivalent organobismuth compounds can be oxidized by Chloramine T to Bi(V) (Scheme 26).107 Fe(IV) complexes can also prepared using this reagent,108 or by the related hypervalent iodine reagent N-tosyliminobenzyliodinane PhdI]NdTs.109

Scheme 26 Generation of a Bi(V) imide by using chloramine T.

The direct oxidation of metal centers with N2 to generate high valent metal complexes is in principle possible, but difficult due to the low reactivity of N2 under ambient conditions. However, N2 can be activated by some metal complexes, the most famous example (although not organometallic) being Mo(III) amido complexes that react to give Mo(VI) nitrido complexes (Scheme 27).110 Interestingly, trimesitylmolybdenum Mo(Mes)3 is also reactive towards N2, resulting in a dimeric Mo(V) complex.111 The corresponding mononuclear Mo(VI) nitrido complex (Mes)3Mo^N is also known.112

Scheme 27 Reactions of N2 with Mo(III) complexes.

For the preparation of high valent metal nitrido complexes, azido complexes are much better starting materials, since they can release N2 (formally a two-electron oxidation of the metal). Since heating of metal azides should be avoided, irradiation with light can trigger the decomposition of the azido ligand. A typical unwanted side reaction is photoreduction, where one electron is transferred from N−3 to the metal ion, which leaves a reduced metal ion (Scheme 28).113 Specific examples for the generation of highvalent metal complexes prepared by this route will be discussed in the carbene section, which are used as co-ligands.

Scheme 28 Photochemistry of metal azides to generate high-valent nitrido complexes.

Very High Oxidation States in Organometallic Chemistry

121

Another ingenious way to generate high valent metal nitrido complexes is the substitution of a chloride ion in a Fe(II) complex by the anionic ligand Dbabh (dbabh: 2,3:5,6-dibenzo-7-aza bicyclo[2.2.1]hepta-2,5-diene). The newly generated Fe(II) amide complex decomposes to give a Fe(IV) nitride (Scheme 29). The formation of aromatic anthracene is the driving force of the reaction. Interestingly, the Fe(IV) nitride is unstable in vacuum and forms a dimeric Fe(I)dN2dFe(I) complex.114 Despite the great synthetic potential of the Dbabh ligand, it has been rarely used in organometallic chemistry.115

Scheme 29 Preparation of an Fe(IV) nitride.

High-valent nitrido complexes can also be reacted with transmetallating agents. One especially remarkable example is the generation of Ru(VI) alkyl complexes. Transmetalation of [Ru(N)Cl4]− or [Ru(N)(OSiMe3)4]− with organomagnesium or organoaluminium reagents leads to the corresponding Ru(VI) nitrido complexes with four alkyl groups [Ru(N)R4]− (Scheme 30).116 These thermally stable compounds are sensitive to oxygen and light, though. The corresponding osmium complexes have also been obtained.117 By treatment with the acid HBF4 in presence of additional ligands L (for example pyridine), neutral or cationic Ru(VI) nitride complexes can be isolated, e.g. [Ru(N)R3L] and [Ru(N)R2L2]+ BF−4.118 The corresponding ruthenium(VI) aryl complexes seem to have not been investigated.

Scheme 30 Preparation of Ru(VI) alkyl complexes with nitrido ligands.

In principle, high valent metal nitrido complexes can also be generated by metathesis reactions without change of oxidation states. Metal carbyne complexes are reacted with nitriles to give nitrido complexes and alkynes (Scheme 31),119 but this approach lacks generality.

Scheme 31 Preparation of an organometallic W(VI) nitride complex by metathesis.

1.05.5

Carbenes and carbynes

During the past decades, carbenes have become one of the most important ligand classes in organometallic chemistry. Differences, similarities and applications between the different classes of carbene ligands (e.g. Fischer, Schrock and N-heterocyclic carbenes) have been reviewed elsewhere.120 As a consequence of the attempts to prepare high valent alkyl complexes of early transition metals,

122

Very High Oxidation States in Organometallic Chemistry

the so-called “Schrock carbenes” were discovered. For example, homoleptic Ta(V) alkyls (without b-hydrogens) can undergo intramolecular a-hydride abstraction (Scheme 32; here: elimination of neopentane).121 Upon treatment with strong bases these carbene complexes can be further deprotonated to give carbyne complexes (Scheme 32).122 In contrast to their corresponding Fischer analogues, Schrock carbenes/carbynes contain typically early transition metals in high oxidation states. In terms of formal oxidation states, Schrock carbenes are regarded as dianionic ligands, while the corresponding carbynes are trianionic and also strong p-donor ligands.

Scheme 32 Preparation of Schrock carbenes and carbynes.

Similarly, carbynes of tungsten in oxidation state +6 can be prepared, although the yields depend significantly on the choice of the starting material and the transmetallating agent (Scheme 33).123

Scheme 33 Preparation of a W(VI) carbyne.

A similar procedure has not yet led to carbynes of rhenium in oxidation state +7. However, the reaction of ReCl4(THF)2 with Me3SiCH2MgCl under a nitrogen atmosphere gives a complicated mixture of compounds. In addition to the bridging dinitrogen complex [(Me3SiCH2)4Re]2N2 or bridging alkylidyne complex Re2(m2-CSiMe3)2(CH2SiMe3)4, the monomeric Re(VII) carbyne complex ReCl(CH2SiMe3)3(CSiMe3) is isolated in 10% yield (Scheme 34).124 This reaction demonstrates the complexity of reactions between transition metal halides and organometallic reagents. Despite the reducing character of the organomagnesium reagent, the oxidation state of rhenium increases from +4 to +7, from an apparent disproportionation pathway.

Scheme 34 A rhenium(VII) carbyne complex.

In the past decades, N-heterocyclic carbenes (NHC) as strong s-donor ligands have evolved into an extremely important ligand class in organometallic chemistry. Recently, their suitability to stabilize high oxidation states of 3d metals has been reviewed.125 For electrostatic reasons, metals in high oxidation states need anionic ligands to balance the charges. As NHC are neutral ligands (and additionally for steric reasons), typically only a small number of NHC ligands are coordinated to a metal in a high oxidation state. A fascinating example how the proper choice of powerful s and p-donor ligands can stabilize high formal oxidation states is the isolation of a Co(V)-diimido-carbene complex.126 Its synthesis involves the reaction of a Co(0)-NHC complex with two equivalents of the organic azide DippN3 (Dipp ¼ 2,6-diisopropylphenyl). The labile olefin ligand and two equivalents of nitrogen are liberated,

Very High Oxidation States in Organometallic Chemistry

123

resulting in the formation of a Co(IV)-NHC-diimido complex. Subsequent one-electron oxidation with ferrocenium salts gives the corresponding Co(V) complex (Scheme 35). For comparison, the oxidation state +5 for cobalt cannot even be reached in purely inorganic compounds with elemental fluorine at high pressure/high temperatures (formation of Cs2CoF6).127

Scheme 35 Preparation of a Co(V) NHC complex.

Another spectacular example is the isolation of formally Ni(IV) NHC complexes prepared by the oxidative halogenation of the corresponding Ni(II) complexes. (Scheme 36).128

Scheme 36 A stable Ni(IV) NHC complex.

Other notable high valent NHC complexes are nitrido complexes of Fe(IV) and Fe(V).129 They are prepared by photolysis of Fe(II) azido complexes, which eliminate N2. Then, the Fe(IV) nitrido complex can be chemically oxidized with ferrocenium to the corresponding Fe(V) nitrido complex (Scheme 37).

Scheme 37 Preparation of a Fe(V) nitrido carbene complex.

Regarding the highest possible oxidation state of metal carbene complexes, rhenium is the metal of choice. Although Re2O7 reacts instantaneously with free NHCs (accompanied by reduction), CH3ReO3 however forms a stable adduct with two equivalents of 1,3-dimethylimidazolin-2-ylidene.130 Interestingly, Re2O7 is cleaved by carbodiphosphorane Ph3P]C]PPh3 to give a cationic organorhenium(VII) complex [(Ph3P)2CReO3]+ ReO−4.131 Due to the air sensitivity of many free N-heterocyclic carbenes, reactions with strongly oxidizing metal complexes can be problematic. For instance, the reaction of the powerful fluorinating agents AuF3 and [AuF4]− to the N-heterocyclic carbene SIMes gives [AuF3(SIMes)], when dichloromethane is used as a solvent at low temperatures (Scheme 38). However, dichloromethane reacts even at low temperatures with AuF3 and [AuF4]−. Interestingly, fluorination of the NHC was observed as a side reaction only when [AuF4]− was used as starting material, but not with the more reactive AuF3.132

124

Very High Oxidation States in Organometallic Chemistry

Scheme 38 Preparation of organogold(III) fluorides using the N-heterocyclic carbene SIMes.

One remarkable NHC with regard to oxidative stability is 1,3-dimesityl-4,5-dichloroimidazol-2-ylidene. This NHC with an unsaturated and chlorinated backbone is quantitatively formed by reaction of commercially available 1,3dimesitylimidazol-2-ylidene (IMes) with CCl4 at room temperature (Scheme 39).133 The chlorinated carbene ligand is far less sensitive than the non-chlorinated one, for example towards oxygen. Thus, the chlorinated carbene forms stable complexes even with extremely aggressive reagents such as SbF5 or AsF5 using 1,3-bis(trifluoromethyl)benzene as oxidation-stable solvent!134 These results suggest that this carbene might be a privileged ligand for highly reactive, high oxidation state complexes of transition metals in the future.

Scheme 39 Preparation of chlorinated NHC ligands and their reactivity towards strong Lewis acids.

1.05.6

Cyclopentadienyl complexes

The stability of high-valent metal-cyclopentadienyl complexes is highly dependent on the substitution patterns of the cyclopentadienyl ring, which influence the sterics and electronics. Decamethylferrocene Cp 2Fe is more electron-rich than ferrocene Cp2Fe (oxidation potential in CH3CN: −0.59 V, in CH2Cl2: −0.48 V vs Cp2Fe).135 As a consequence, both electrochemical and chemical oxidation to [Cp 2Fe]2+ are possible (oxidation potential +1.31 V vs AgRE in liquid SO2).136 Chemical oxidation to the dication occurs with strong fluorooxidizers like SbF5, AsF5 or ReF6 in SO2 or XeF+ Sb2F−11 in anhydrous HF and salts of the [Cp 2Fe]2+ dications can be isolated and structurally characterized.137 Depending on the weakly-coordinating counteranion, the Cp rings are either parallel (in case of the Sb2F−11 salt) or slightly tilted (in the SbF−6 salt—Scheme 40). In contrast, electrochemical oxidation of ferrocene Cp2Fe occurs at much more positive potentials (+1.79 V vs AgRE in liquid SO2) and generates a highly unstable dication.136 Interestingly, Cp 2Fe2+ forms a stable Fe(IV) carbonyl complex. Although one could expect that such an Fe(IV) species would fall into the category of non-classical carbonyl complexes (meaning that ṽ(CO) > 2143 cm−1),138 the ṽ(CO) of 2034 cm−1 in [Cp 2FeCO][SbF6]2 indicates that backbonding is taking place, which is surprising for such a high oxidation state but is a consequence of ligand-to-ligand charge transfer (Cp to CO).139

Scheme 40 Molecular structure of Cp 2Fe2+ in the crystal structure of [Cp 2Fe][SbF6]2 2HF.

Very High Oxidation States in Organometallic Chemistry

125

Dications of other permethylated metallocenes can also be isolated. As a 20-valence electron complex Cp 2Ni is easily oxidized to the formally nickel(IV) complex [Cp 2Ni]2+ by weak oxidants like Br2 or silver salts.140 Decamethylmanganocene Cp 2Mn is oxidized to [Cp 2Mn]2+ [SbF6]−2 by SbF5 in anhydrous HF. The oxidation of highly stable Cp 2Co+ SbF−6 to the 17-electron Co(IV) complex [Cp 2Co]2+ is only achieved by extreme oxidants like IrF6 or O+2 Sb2F−11 in anhydrous HF/SbF5.141 Additionally, the crystal structure of [Cp 2Pt]2+ [PF6]−2 is known.142 By introducing additional strong p-donor ligands (e.g. oxo) it is possible to access even higher oxidation states. The preparation of the pentamethylcyclopentadienyl complex of rhenium(VII), Cp ReO3, has been reported by Herrmann.143 Irradiation of the Re(I) precursor Cp Re(CO)3 leads to the substitution of one carbonyl ligand against THF. Afterwards, air oxidation leads to the corresponding trioxorhenium(VII) complex (Scheme 41). Analogous CpReO3 is so far unknown. In general, by substituting Cp by Cp the ionization energies of the corresponding molecules decrease, resulting in more facile oxidation.144

Scheme 41 Preparation of a Re(VII) pentamethylcyclopentadienyl complex.

1.05.7

Hydride

Due to its small size, anionic charge and good s-donor properties, hydrides have the potential to stabilize high formal oxidation states (e.g. ReH2− 9 ). There are various ways to prepare high valent metal complexes with hydride ligands, the most important being reactions with dihydrogen H2, reactions with hydride sources and protonation reactions. However, the potential for equilibrium between metal dihydrido and metal dihydrogen complexes complicates the picture, since the reductive elimination of the hydrides formally lowers the oxidation state by two units (Scheme 42). In addition, dihydrogen complexes have variable HdH distances that suggest that complexes can lie between these extremes (non-classical hydrogen complexes). The delicate balance between both extremes depends on the choice of the metal, other ligands etc.

Scheme 42 Metal-dihydrido vs metal-dihydrogen complexes.

Oxidative addition of H2 to low-valent metal centers is a well-established reaction in organometallic chemistry. For example, the formally Fe(IV)-dihydrido complex [Cp Fe(dippe)H2]+ [BPh4]− (dippe ¼ 1,2-bis(diisopropylphosphino)ethane) can be prepared from an electron-rich Fe(II) precursor by reaction with H2 (Scheme 43).145 Nevertheless, oxidative addition of H2 becomes more and more difficult with increasing oxidation state of the metal.

Scheme 43 Preparation of an Fe(IV) dihydrido complex.

Protonation reactions of organometallic compounds are another route to high-valent hydride complexes. By definition, a single hydrogen bound to a metal is considered as a hydride H−. Consequently, if a metal center is protonated, its oxidation state formally increases by +2. As a recent example, the protonation of ferrocene Cp2Fe in the superacidic mixture HF/PF5 gives the cationic formally Fe(IV) hydrido complex [Cp2FeH]+ PF−6 (Scheme 44).146

126

Very High Oxidation States in Organometallic Chemistry

Scheme 44 Protonation of ferrocene.

However, protonation of metal hydride complexes does not necessarily further increase the oxidation state of the metal, if it is followed by reductive elimination to give dihydrogen. The reaction of the Ru(IV) trihydrido complex Cp RuH3(EPh3) (E ¼ Sb, As) with HBF4 ∙Et2O gives the non-classical complex [Cp Ru(H2)2(EPh3)]+ BF−4.147 Since there are formally two neutral dihydrogen ligands, the oxidation state of ruthenium is +2 in this product (Scheme 45).

Scheme 45 Formal reduction of Ru(IV) to Ru(II) upon protonation.

In comparison with a monoanionic Z5-cyclopentadienyl ring as a 6-electron donor (ionic counting), three hydride ligands donate the same number of electrons but their in total three negative charges are therefore formally better suited to cancel out higher positive charges of metal ions. A series of mixed cyclopentadienyl-polyhydride complexes nicely illustrates this observation. 18-electron complexes of the composition Cp ReH6 (d0), Cp OsH5 (d2) and Cp IrH4 (d4) are known in the literature. These very high formal oxidation states could probably not be achieved in the corresponding dicyclopentadienyl complexes without any further anionic ligands. However, as a consequence of the highly covalent nature of the MdH bond, the metal centers in the abovementioned examples are not necessarily very oxidized (in terms of physical oxidation state), which is also reflected by their synthesis. Pentagonal-bipyramidal Cp ReH6 is obtained by reacting the Re(V) precursor Cp ReOCl2 with LiAlH4, followed by careful quenching with methanol (Scheme 46).148 The absence of Z2-H2 ligands was confirmed later via gas phase electron diffraction and single crystal X-ray diffraction (the latter for (C5Me4Et)ReH6, confirming the oxidation state of +7 for rhenium.149

Scheme 46 Preparation of Cp ReH6.

A similar reaction gives access to a pentagonal-pyramidal Os(VI) complex, by treatment of the Os(III) dimer [Cp OsBr2]2 with LiAlH4, followed by reaction with methanol (Scheme 47).150

Scheme 47 Preparation of Cp OsH5.

Cp IrH4, a formally iridium(V) complex, is prepared in a two-step process.151 First, dimeric [Cp IrCl2]2 is reacted with dihydrogen and ammonium hexafluorophosphate to give a dimeric, cationic Ir(III) complex with three m2-bridging hydride ligands. Treatment with the “reducing agent” lithium triethylborohydride gives the formally oxidized species Cp IrH4 (Scheme 48). Interestingly, the corresponding Ir(V) alkyl complex is also known. Cp Ir(CH3)4 can also be prepared from [Cp IrCl2]2 and trimetylaluminium Al2(CH3)6 in 11% yield.152

Very High Oxidation States in Organometallic Chemistry

127

Scheme 48 Preparation of Cp IrH4.

In summary, the hydride ligand seems to be a rather overlooked ligand to stabilize high oxidation states, especially when compared for example with the ubiquitous oxo ligand. However, especially in case of polyhydride complexes, the non-trivial spectroscopic differentiation between hydride and dihydrogen ligands complicates the assignment of both structure and oxidation states.153

1.05.8

Silyl chemistry

Like the hydride ligand, silyl groups R3Si are good s-donor ligands and can effectively stabilize high oxidation states. Especially valuable for the synthesis of such complexes is the tendency of SidH bonds to undergo oxidative addition to metal centers. Brookhart investigated the oxidative addition of diphenyldisilane Ph2SiH2 to Co(I) complexes, yielding dihydridodisilylcobalt(V) complexes Cp Co(H)2(SiR3)2 (Scheme 49).154 Similarly, reaction of three-coordinate Co(I) pincer complexes with an excess of PhSiH3 produces seven-coordinate Co(V) complexes with hydride and silyl ligands.155

Scheme 49 Preparation of a Co(V) complex.

The oxidative addition of silanes, e.g. Et3SiH to the Rh(III) complex [Cp RhCl2]2 permits the isolation of the Rh(V) complex Cp Rh(H)2(SiEt3)2 (Scheme 50).156 The corresponding iridium complex has been prepared by the same synthetic route.157 

Scheme 50 Preparation of a Rh(V) complex.

However, the oxidative addition of SidH bonds can be incomplete, because it is possible for SidH bonds to interact with the metal much like in a dihydrogen complex, without full oxidative addition. For example, addition of bis(2-silylphenyl)silane to a Ni(0) nickel complex has been reported to give a complex that can exist either as a Ni(II) bis(silyl)[Z2-Si-H] complex or as a Ni(IV) trisilyl hydride complex (Scheme 51). In the solid state, a crystal structure supports the presence of the former, while in solution at low temperatures the latter species is present (detection via NMR spectroscopy).158

Scheme 51 Agostic interactions in a silylhydrido complex of nickel.

128

Very High Oxidation States in Organometallic Chemistry

Another very successful strategy to prepare high valent metal silyl complexes is the use of 1,2-disilylbenzene derivatives. Oxidative addition of two equivalents of 1,2-disilylbenzene with the Ni(0) compound Ni(dmpe)2 (dmpe ¼ 1,2-bis(dimethylphosphino)ethane) gives a tetrakis(silyl) Ni(IV) complex in 81% yield after 10 days of heating at 80  C in toluene (Scheme 52).159 Similar tetrasilyl Pd(IV) and Pt(IV) complexes are also known.160

Scheme 52 Preparation of a tetrakis(silyl) Ni(IV) complex.

In this context the preparation of a formally hexavalent hexasilylpalladium complex has to be mentioned (Scheme 53).161 This finding was discussed intensively in the literature since the alternative description as a Pd(II) complex with two Z2-disilane ligands seems to better suited.162 The ambiguity between these forms reminds of the differentiation between polyhydrides and non-classical dihydrogen complexes.

Scheme 53 Alternative bonding descriptions for a hexasilyl palladium complex.

The fact that metal complexes in unusually high formal oxidation states can be stabilized by either hydride or silyl ligands might be explained by their comparatively low electronegativity, which makes the metal-ligand bond very covalent, and the (ionic) oxidation state formalism unrealistic.155

1.05.8.1

Halogenation reactions

Oxidative addition of elemental halogens X2 or the corresponding acids HX to low-valent metal centers is a well known reaction.31 Due to the stronger oxidizing character of the lighter halogens, they are better suited to access high oxidation states of metal centers. Additionally, from an electrostatic perspective, small anionic p-donor ligands like F− are optimal for the stabilization of high-valent metal ions. However, due to the higher reactivity of Cl2 and F2 (especially via radical pathways which may lead to the halogenation of ligand CdH bonds) and the difficulties in handling these corrosive and toxic gases, the investigation of such reactions is challenging. Furthermore, reductive elimination from the resulting high-valent complexes can occur. Despite four electron-withdrawing trifluoromethyl groups, Pt(IV) complexes are accessible by direct halogenation of Pt(II) complexes (Scheme 54).163

Scheme 54 Oxidative addition of halogens to Pt(II) complexes.

Very High Oxidation States in Organometallic Chemistry

129

Instead of elemental halogens, it is also possible to oxidize complexes using commercially available XeF2, which is a solid, powerful oxidizing and fluorinating agent. Although the use of perfluorinated polymers (Teflon, PFA) for reaction vessels is recommended, reactions can also be performed in glass, however the vessel material may have an influence on the reactivity of XeF2 in some reactions.164 Typically, XeF2 is best used in oxidation-stable solvents as acetonitrile. Dichloromethane reacts slowly enough with XeF2 to be a useful solvent. In contrast, chloroform is fluorinated in significant amounts by XeF2.165 For example, XeF2 in dichloromethane at 0  C cleanly fluorinates the organogold(I) species [Au(CF3)2]− to the corresponding gold(III) complex [(CF3)2AuF2]− (Scheme 55).166

Scheme 55 Preparation of Au(III) complexes with XeF2.

Also chloro- and 1,2-dichlorobenzene can be used as suitable solvents for fluorinations with XeF2, although the latter has to be used slightly in excess. This has been demonstrated in the preparation of Ni(III) and Ni(IV) mixed fluoride and trifluoromethyl complexes (Scheme 56).167

Scheme 56 Preparation of Ni(III) and Ni(IV) fluorides with XeF2.

Recent uses of XeF2 in organometallic chemistry include the fluorination of arylboronic ester catalyzed by organobismuth redox catalysis. The oxidation power of XeF2 is sufficient to oxidize Bi(III) to Bi(V) fluoride complexes during the catalytic cycle.168 Electrophilic fluorination reagents (typically with NdF bonds, e.g. Selectfluor) formally transfer “F+” to substrates.169 Many of these easily handable compounds are commercially available. Their relative reactivities and their electrochemical properties are well known.170As an example, oxidative fluorination of organometallic Ni(II) complexes with Selectfluor gives access to Ni(IV) fluoride species. Upon heating, these complexes can decompose via reductive elimination and formation of CdF bonds (Scheme 57).171

Scheme 57 Electrophilic fluorination of Ni(II).

1.05.9

One-electron oxidizing agents

Connelly and Geiger have reviewed in detail the literature on various reducing and oxidizing agents and the relationship to electrochemistry.135 In general, the oxidation of substrates with oxidation potentials of >+1 V vs Cp2Fe+/Cp2Fe is challenging since only a few oxidants are sufficiently strong but have sufficient selectivity as a simple one-electron oxidant. For example, halogens X2

130

Very High Oxidation States in Organometallic Chemistry

are strong oxidants, but often cause M-X bond-forming reactions like oxidative addition. Additionally, the choice of a suitable solvent is of utmost importance. For the isolation of a highly oxidized metal complex, the solvent must be oxidation-stable and, for unstable species, must have a low melting point to allow crystallization without warming. Acetonitrile is often an excellent solvent choice, and propionitrile shows similar stability with a lower melting point. Beside dichloromethane also halogenated arenes (e.g. ortho-dichloro- and ortho-difluorobenzene) have been used lately as weaklycoordinating and oxidation stable solvents but the melting points of the halogenated arenes are also relatively high. However, they react with some oxidizers (like NO+ and NO+2).172 Although sulfur dioxide is a toxic gas at ambient temperature, it is an outstanding solvent for electrochemical investigations at very positive potentials (electrochemical window −1.1 V to +3.3 V vs Cp2Fe+/Cp2Fe).173 Due to its sufficiently broad liquid range of −75  C to −10  C it can be conveniently used as a solvent, however its vapor pressure of about 3 bars at 20  C requires special attention during manipulations. SO2ClF has a similar oxidative stability, but a wider liquid range (−125  C to +7  C) and therefore a lower vapor pressure at room temperature. The most oxidation stable solvent is anhydrous hydrogen fluoride HF. Since its rapidly destroys glass and is highly toxic, it can only be handled in a stainlesssteel vacuum line.174 For example, it proved to be the solvent of choice for the isolation and crystallization of the Fe(IV) species Cp 2Fe2+.137 The suitability of ferrocenium as an oxidant to generate some high-valent metal centers has been mentioned already in this chapter. In addition, electron-withdrawing groups increase the oxidation power of a substituted ferrocene.135 However, ferrocenium is comparatively not a very powerful oxidant. Nitrosyl salts (containing NO+) have been used for decades as oxidation agents. The oxidation potential of +1 V vs Cp2Fe+/Cp2Fe in dichloromethane is probably slightly underestimated.135 Nitrosyl salts offer the advantage that many of them are commercially available. NO gas is released and can be removed from the reaction vessel in vacuum. Nitryl salts NO+2 are even slightly stronger oxidizing than NO+ salts (+2.1 V vs NHE for NO+2 and +1.7 V vs NHE for NO+), however they are available with fewer anions.175 Poleschner and Seppelt have developed a system (Scheme 58) where commercially available XeF2 is used in combination with fluoride acceptors to oxidize organic or organometallic compounds (LA ¼ BF3, B(C6F5)3, Al[(OC(CF3)3]3; R3Si-A ¼ Me3SiOTf, + − Me3SiNTf2, Me3SiOTeF5, Me3Si+ B(C6F5)−4, (Et3Si)+2 B12Cl2− 12 and Me3Si CHC11Cl11). This system allows the introduction of various weakly-coordinating anions to increase the chance of crystallizing an oxidized substrate. The reactions can be performed in regular glassware in water-free solvents like CH2Cl2, C2H5CN, ortho-dichlorobenzene, and liquid SO2. The oxidation limit of this system has been estimated to be  +2 V vs NHE.176

Scheme 58 Introduction of weakly-coordinating anions during oxidations with XeF2.

Triphenylaminium radical ions are well established one-electron oxidants, the most prominent being “Magic Blue”—tris(4bromophenyl)aminium hexachloroantimonate. This commercially available compound displays a relatively high redox potential of +0.67 V in acetonitrile (vs Cp2Fe+/Cp2Fe).135 By further increasing the degree of halogenation, the redox potentials are shifted to higher values. The corresponding tris(2,4-dibromophenyl)aminium and tris(2,4,6-tribromophenyl)aminium salts display redox potentials of 1.14 V and 1.36 V.177 Similarly, the redox potential of the aminium radical cation of hexabrominated phenylcarbazole is 1.34 V (in CH2Cl2, vs Cp2Fe+/Cp2Fe) which is strong enough to oxidize fullerene C60 to its radical cation (Scheme 59).178 Recently, Krossing reported the use of the radical cation of a perfluorinated dihydrophenazine as a powerful oxidant (+1.21 V vs Cp2Fe+/Cp2Fe) in MeCN and at 1.29 V in ortho-difluorobenzene. In presence of other ligands L as CO or tBu-NC, this oxidant is able to oxidize decamethylferrocene Cp 2Fe to the corresponding Fe(IV) complexes [Cp 2Fe(L)]2+.179

Scheme 59 Nitrogen-based radical cations as powerful one-electron oxidizers.

Very High Oxidation States in Organometallic Chemistry

131

The oxidation of the neutral amines to their radical cations is typically achieved by the use of SbCl5, SbF5 or AsF5.135 Alternatively, soluble silver salts with weakly coordinating anions are used in combination with elemental halogens X2 (X ¼ Cl, Br, I). The polarization of the X2 molecule by the coordination of Ag+ as well as the precipitation of insoluble silver halides AgX explain the higher oxidation power of these systems compared to the free halogens.180 Fluorinated oxidizers like SbF5 and AsF5 are both Lewis acids and oxidants (Scheme 60). As oxidants they offer the advantage of generating stable hexafluorometallates as counteranions, and the corresponding trifluoride SbF3 and AsF3 are formed as byproducts. While AsF3 is a volatile (although toxic) liquid that can be easily pumped off, polymeric SbF3 is a non-volatile solid. SbF5 is a highly viscous and corrosive liquid which can in principle be handled inside a glovebox. Since weighing of small amounts of SbF5 poses a significant problem, the use of the crystalline adduct SbF5 ∙SO2 is recommended.181 AsF5 is a highly toxic and corrosive gas with a low boiling point (−53  C). It should only be handled in stainless-steel equipment by trained personnel. Due to the extreme lewis acidity of SbF5 and AsF5 they are usually not used in coordinating solvents (due to adduct formation)182 while Cl/F exchange or other side reactions could happen in chlorinated solvents.183 For oxidation purposes the use of SO2 or SO2ClF as oxidation-stable solvents is optimal. AsF5 is a much stronger oxidant than SbF5. While the oxidation power of SbF5 is just high enough to quantitatively oxidize decametylferrocene Cp 2Fe to its dication after several minutes at room temperature (redox potential of Cp 2Fe2+/Cp 2Fe+ in SO2: +1.23 V vs Cp2Fe+/Cp2Fe), the same reaction occurs much more rapidly and at lower temperatures with AsF5. AsF5, but not SbF5, has been demonstrated to be able to oxidize the perhalogenated boron cluster B12Cl2− 12 to its monoanion and even its neutral form in liquid sulfur dioxide which occur at potentials of +2.11 V and +2.67 V vs Cp2Fe+/Cp2Fe!173

Scheme 60 Oxidizing properties of AsF5 and SbF5.

The use of metal hexafluorides as one-electron oxidizers has been very limited in organometallic chemistry, but has potential for the future. The electron affinities of MF6 (to give MF−6) vary considerably depending of the choice of the metal, and their oxidation power increases in the series WF6 < MoF6 < ReF6 < OsF6 < IrF6 < PtF6 (other 4d metal hexafluorides beside Mo are very difficult to prepare).184 The redox potential of WF6 is about +1.08 V vs SCE), which is significantly lower than of WCl6 (+1.59 V vs SCE).185 MoF6 is even stronger (+2.08 V vs SCE). Due to the even higher oxidation power of ReF6 and OsF6 these hexafluorides are usually not used in organic solvents, but handled in liquid SO2. ReF6 in SO2 is able to oxidize decamethylferrocene Cp 2Fe to its dication.137 IrF6 and PtF6 are such powerful oxidizing and fluorinating agents that they can only be handled in anhydrous HF. In SO2, formation of SO2F2 occurs. PtF6 is known for its ability to oxidize elemental xenon or dioxygen directly at room temperature.186 The corresponding dioxygenyl salts (containing O+2) are also very powerful oxidants. Their oxidation potential has been estimated to be +5 V vs NHE.187 O+2 Sb2F−11 as well as IrF6 are strong enough to oxidize Cp 2Co+ to Cp 2Co2+ in HF/ SbF5.141

1.05.10

Conclusion

The organometallic chemistry of metal complexes in very high oxidation states has a long history, and is continuing to develop. Seminal works of Wilkinson and others blazed the way to exciting compounds that enrich our understanding of organometallic chemistry. But still today, as spectroscopic methods are far more developed, still many open questions remain and many possibilities for further studies exist. One intention of this review has been to demonstrate that ligand design in combination with appropriate co-ligands opens the way to organometallic complexes with unusually high formal oxidation states. Surprisingly, many examples of these very high oxidation state complexes can use common solvents and do not require strongly oxidizing conditions, revealing the limitations of the oxidation-state formalism. However, these complexes continue to push the borders of organometallic chemistry. There are grounds for optimism that for some systems the application of fluorine-based oxidants in solvents such as SO2 or HF could give access to additional highly oxidized species in the future.

Acknowledgement M.M. acknowledges funding by Freie Universität Berlin.

References 1. 2. 3. 4.

Cotton, F. A.; Wilkinson, G.; Murillo, C. A.; Bochmann, M. Advanced Inorganic Chemistry, 6th ed.; John Wiley & Sons, Inc., 1999 Eisch, J. J.; Adeosun, A. A.; Dutta, S.; Fregene, P. O. Eur. J. Inorg. Chem. 2005, 2657–2670. Simoes, J. A. M.; Beauchamp, J. L. Chem. Rev. 1990, 90, 629–688. Morrison, J. A. Adv. Organometal. Chem. 1993, 35, 211–239.

132

Very High Oxidation States in Organometallic Chemistry

5. (a) Bower, B. K.; Tennent, H. G. J. Am. Chem. Soc. 1972, 94, 2512–2514; (b) Bower, B. K. Patent 3705916, 12.12.1972. 6. (a) Byrne, E. K.; Richeson, D. S.; Theopold, K. H. J. Chem. Soc. Chem. Commun. 1986, 1491–1492; (b) Byrne, E. K.; Theopold, K. H. J. Am. Chem. Soc. 1987, 109, 1282–1283; (c) Byrne, E. K.; Theopold, K. H. J. Am. Chem. Soc. 1989, 111, 3887–3896. 7. Lewis, R. A.; Smiles, D. E.; Darmon, J. M.; Stieber, S. C. E.; Wu, G.; Hayton, T. W. Inorg. Chem. 2013, 52, 8218–8227. 8. (a) Dimitrov, V.; Linden, A. Angew. Chem. Int. Ed. 2003, 42, 2631–2633; (b) Liptrot, D. J.; Guo, J.-D.; Nagase, S.; Power, P. P. Angew. Chem. Int. Ed. 2016, 55, 14766–14769. 9. Carnes, M.; Buccella, D.; Chen, J. Y.-C.; Ramirez, A. P.; Turro, N. J.; Nuckolls, C.; Steigerwald, M. Angew. Chem. Int. Ed. 2009, 48, 290–294. 10. (a) Li, H.; Hu, Y.; Zhang, Z.; Fan, Q.; King, R. B.; Schaefer, H. F., I J. Phys. Chem. A 2019, 123, 9514–9519; (b) Li, H.; Wang, L.; Hu, Y.; Zhang, Z.; Wan, D.; Fan, Q.; King, R. B.; Schaefer, H. F., I J. Phys. Chem. A 2020, 124, 6867–6876. 11. Casitas, A.; Rees, J. A.; Goddard, R.; Bill, E.; DeBeer, S.; Fürstner, A. Angew. Chem. Int. Ed. 2017, 56, 10108–10113. 12. Kolodziej, R. M.; Schrock, R. R.; Davis, W. M. Inorg. Chem. 1988, 27, 3253–3255. 13. Taullaj, F.; Armstrong, D.; Datta, S.; Lough, A. J.; Fekl, U. Eur. J. Inorg. Chem. 2019, 1288–1291. 14. Bochmann, M.; Wilkinson, G.; Young, G. B. J. Chem. Soc. Dalton. Trans. 1980, 1879–1887. 15. Claire, P. P. K.; Jones, C. J.; McCleverty, J. A.; Coe, P. L.; Drew, M. G. B. J. Organomet. Chem. 1992, 424, 105–114. 16. Shortland, A. J.; Wilkinson, G. J. Chem. Soc. Dalton. Trans. 1973, 872–876. 17. (a) Galyer, A. L.; Wilkinson, G. J. Chem. Soc. Dalton. Trans. 1976, 2235–2238; (b) Galyer, L.; Mertis, K.; Wilkinson, G. J. Organomet. Chem. 1975, 85, C37–C38. 18. Pfennig, V.; Seppelt, K. Science 1996, 271, 626–628. 19. (a) Kang, S. K.; Albright, T. A.; Eisenstein, O. Inorg. Chem. 1989, 28, 1611–1613; (b) Landis, C. R.; Cleveland, T.; Firman, T. K. J. Am. Chem. Soc. 1995, 117, 1859–1860. 20. Roessler, B.; Seppelt, K. Angew. Chem. Int. Ed. 2000, 39, 1259–1261. 21. Mertis, K.; Wilkinson, G. J. Chem. Soc. Dalton. Trans. 1976, 1488–1492. 22. Pfennig, V.; Robertson, N.; Seppelt, K. Angew. Chem. Int. Ed. 1997, 36, 1350–1352. 23. Gilman, H.; Woods, L. A. J. Am. Chem. Soc. 1948, 70, 550–552. 24. (a) Zhu, D.; Lindeman, S. V.; Kochi, J. K. Organometallics 1999, 18, 2241–2248; (b) Rice, G. W.; Tobias, R. S. Inorg. Chem. 1975, 14, 2402–2407. 25. (a) Naumann, D.; Roy, T.; Tebbe, K.-F.; Crump, W. Angew. Chem. Int. Ed. 1993, 32, 1482–1483; (b) Willert-Porada, M. A.; Burton, D. J.; Baenziger, N. C. J. Chem. Soc. Chem. Commun. 1989, 1633–1634; (c) Dukat, W.; Naumann, D. Rev. Chim. Miner. 1986, 23, 589–603; (d) Schlueter, J. A.; Williams, J. M.; Geiser, U.; Dudek, J. D.; Sirchio, S. A.; Kelly, M. E.; Gregar, J. S.; Kwok, W. H.; Fendrich, J. A.; Schirber, J. E.; Bayless, W. R.; Naumann, D.; Roy, T. J. Chem. Soc. Chem. Commun. 1995, 1311–1312. 26. Baya, M.; Joven-Sancho, D.; Alonso, P. J.; Orduna, J.; Menjón, B. Angew. Chem. Int. Ed. 2019, 58, 9954–9958. 27. DiMucci, I. M.; Lukens, J. T.; Chatterjee, S.; Carsch, K. M.; Titus, C. J.; Lee, S. J.; Nordlund, D.; Betley, T. A.; MacMillan, S. N.; Lancaster, K. M. J. Am. Chem. Soc. 2019, 141, 18508–18520. 28. Guerra, M. A.; Bierschenk, T. R.; Lagow, R. J. J. Organomet. Chem. 1986, 307, C58–C62. 29. Cotton, S. A. Silver and Gold. In Chemistry of the Precious Metals, Dordrecht: Springer, 1997; pp 273–327. 30. Joven-Sancho, D.; Baya, M.; Martin, A.; Menjón, B. Chem. Eur. J. 2018, 24, 13098–13101. 31. Halpern, J. Acc. Chem. Res. 1970, 3, 386–392. 32. Levin, M. D.; Chen, T. Q.; Neubig, M. E.; Hong, C. M.; Theulier, C. A.; Kobylianskii, I. J.; Janabi, M.; O’Neil, J. P.; Toste, F. D. Science 2017, 356, 1272–1276. 33. Nair, H. K.; Morrison, J. A. J. Organomet. Chem. 1989, 376, 149–164. 34. Yagupol’skii, L. M.; Kondratenko, N. V.; Timofeeva, G. N. Zh. Org. Khim. 1984, 20, 115–118. 35. (a) Umemoto, T.; Ishihara, S. Tetrahedron Lett. 1990, 31, 3579–3582; (b) Umemoto, T.; Ishihara, S. J. Am. Chem. Soc. 1993, 115, 2156–2164; (c) Umemoto, T. Chem. Rev. 1996, 96, 1757–1778. 36. Shibata, N.; Matsnev, A.; Cahard, D. Beilstein J. Org. Chem. 2010, 6, 65. 37. Camasso, N. M.; Sanford, M. S. Science 2015, 347, 1218–1220. 38. Siegbahn, P. E. M. J. Phys. Chem. 1995, 99, 12723–12729. 39. Anderson, C. M.; Puddephatt, R. J.; Ferguson, G.; Lough, A. J. J. Chem. Soc., Chem. Commun. 1989, 1297–1298. 40. Xu, L.-M.; Li, B.-J.; Yang, Z.; Shi, Z.-J. Chem. Soc. Rev. 2010, 39, 712–733. 41. Vicente, J.; Arcas, A.; Juliá-Hernández, F.; Bautista, D. Angew. Chem. Int. Ed. 2011, 50, 6896–6899. 42. Bour, J. R.; Camasso, N. M.; Sanford, M. S. J. Am. Chem. Soc. 2015, 137, 8034–8037. 43. Bour, J. R.; Ferguson, D. M.; McClain, E. J.; Kampf, J. W.; Sanford, M. S. J. Am. Chem. Soc. 2019, 141, 8914–8920. 44. Koschmieder, S. U.; Wilkinson, G. Polyhedron 1991, 10, 135–173. 45. Latjaeva, V. N.; razuvaev, G. A.; Malisheva, A. V.; Kiljakova, G. A. J. Organomet. Chem. 1964, 2, 388–397. 46. Ludwig, W.; Seidel, W. Thermochim. Acta 1985, 85, 59–62. 47. Razuvaev, G. A.; Latyaeva, V. N.; Kilyakova, G. A.; Mal’kova, G. Y. Dok. Ak. Nauk. SSSR 1970, 191, 620–621. 48. Wilkinson, G. Sci. Progr. 1993/1994, 77, 15–27. 49. (a) Hay-Motherwell, R. S.; Hussain-Bates, B.; Hursthouse, M. B.; Wilkinson, G. S. J. Chem. Soc. Chem. Commun. 1990, 1242–1243; (b) Hay-Motherwell, R. S.; Wilkinson, G.; Hussain-Bates, B.; Hursthouse, M. B. J. Chem. Soc. Dalton. Trans. 1992, 3477–3482. 50. Hay-Motherwell, R. S.; Wilkinson, G.; Hussain-Bates, B.; Hursthouse, M. B. Polyhedron 1991, 10, 1457–1458. 51. Hay-Motherwell, R. S.; Koschmieder, S. U.; Wilkinson, G.; Hussain-Bates, B.; Hursthouse, M. B. J. Chem. Soc. Dalton. Trans. 1991, 2821–2830. 52. Hay-Motherwell, R. S.; Wilkinson, G.; Hussain-Bates, B.; Hursthouse, M. B. Polyhedron 1993, 12, 2009–2012. 53. Danopoulos, A. A.; Hay-Motherwell, R. S.; Wilkinson, G.; Cafferkey, S. M.; Sweet, T. K. N.; Hursthouse, M. B. J. Chem. Soc. Dalton. Trans. 1997, 3177–3184. 54. So, S.-C.; Cheung, W.-M.; Wang, G.-C.; Huang, E. K.; Lau, M.-K.; Zhang, Q.-F.; Sung, H. H.-Y.; Williams, I. D.; Leung, W.-H. Organometallics 2014, 33, 4497–4502. 55. Arnold, J.; Wilkinson, G.; Hussain, B.; Hursthouse, M. B. J. Chem. Soc. Chem. Commun. 1988, 1349–1350. 56. Arnold, J.; Wilkinson, G.; Hussain, B.; Hursthouse, M. B. J. Chem. Soc. Dalton. Trans. 1989, 2149–2153. 57. Parr, J.; Haiges, R.; Inkpen, M. ChemRxiv; https://doi.org/10.26434/chemrxiv.12830384.v1. 58. Stavropoulos, P.; Savage, P. D.; Tooze, R. P.; Wilkinson, G.; Hussain, B.; Motevalli, M.; Hursthouse, M. B. J. Chem. Soc. Dalton. Trans. 1987, 557–562. 59. Vohs, J. K.; Downs, L. E.; Stasalovich, J.; Barfield, M.; Robinson, G. H. J. Cluster. Sci. 2002, 13, 601–608. 60. (a) Uson, R.; Laguna, A.; Vicente, J. J. Organomet. Chem. 1977, 131, 471–475; (b) Murray, H. H., Jr.; Fackler, J. P.; Porter, L. C.; Briggs, D. A.; Guerra, M. A.; Lagow, R. J. Inorg. Chem. 1987, 26, 357–363. 61. Markwell, A. J. J. Organomet. Chem. 1985, 293, 257–263. 62. Liu, L.; Zhu, M.; Yu, H.-T.; Zhang, W.-X.; Xi, Z. J. Am. Chem. Soc. 2017, 139, 13688–13691. 63. Uson, R.; Forniés, J.; Espinet, P.; Navarro, R.; Martinez, F.; Tomas, M. J. Chem. Soc. Chem. Commun. 1977, 789–790. 64. Alonso, P. J.; Falvello, L. R.; Forniés, J.; Martín, A.; Menjón, B.; Rodríguez, G. Chem. Commun. 1997, 503–504. 65. Alonso, P. J.; Arauzo, A. B.; García-Monforte, M. A.; Martín, A.; Menjón, B.; Rillo, C.; Tomás, M. Chem. Eur. J. 2009, 15, 11020–11030. 66. Fornies, J.; Menjon, B.; Sanz-Carrillo, R. M.; Tomas, M.; Connelly, N. G.; Crossley, J. G.; Orpen, A. G. J. Am. Chem. Soc. 1995, 117, 4295–4304. 67. (a) Usón, R.; Forniés, J.; Tomás, M.; Menjón, B.; Bau, R.; Sünkel, K.; Kuwabara, E. Organometallics 1986, 5, 1576–1581; (b) Usón, R.; Forniés, J.; Tomás, M.; Menjón, B.; Sünkel, K.; Bau, R. J. Chem. Soc. Chem. Commun. 1984, 751–752.

Very High Oxidation States in Organometallic Chemistry 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. 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.

133

Richardson, M. F.; Wulfsberg, G.; Marlow, R.; Zaghonni, S.; McCorkle, D.; Shadid, K.; Gagliardi, J., Jr.; Farris, B. Inorg. Chem. 1993, 32, 1913–1919. García-Monforte, M. A.; Alonso, P. J.; Forniés, J.; Menjón, B. Dalton Trans. 2007, 3347–3359. (a) Debaerdemaeker, T.; Roth, H.; Brune, H.-A. J. Organomet. Chem. 1991, 412, 243–249; (b) Marsh, R. E. Acta Cryst. 1997, B53, 317–322. (a) García, M. P.; Jiménez, M. V.; Lahoz, F. J.; López, J. A.; Oro, L. A. J. Chem. Soc. Dalton. Trans. 1998, 4211–4214; (b) Gardner, S. A.; Gordon, H. B.; Rausch, M. D. J. Organomet. Chem. 1973, 60, 179–188; (c) Suter, D.; van Summeren, L. T. C. G.; Blacque, O.; Venkatesan, K. Inorg. Chem. 2018, 57, 8160–8168; (d) Awad, S. B.; Brown, D. S.; Cohen, S. C.; Humphries, R. E.; Massey, A. G. J. Organomet. Chem. 1977, 127, 127–138; (e) Tugashov, K. I.; Gribanyov, D. A.; Dolgushin, F. M.; Smol’yakov, A. F.; Peregudov, A. S.; Klemenkova, Z. S.; Matvienko, O. V.; Tikhonova, I. A.; Shur, V. B. Organometallics 2017, 36, 2437–2445. Hilton, C. L.; King, B. T. Organometallics 2006, 25, 4058–4061. (a) Ogawa, S.; Sato, S.; Furukawa, N. Tetrahedron Lett. 1992, 33, 7925–7928; (b) Sato, S.; Takahashi, O.; Furukawa, N. Coord. Chem. Rev. 1998, 176, 483–514. Sheppard, W. A. J. Am. Chem. Soc. 1971, 93, 5597–5598. Ogawa, S.; Matsunaga, Y.; Sato, S.; Iida, I.; Furukawa, N. J. Chem. Soc. Chem. Commun. 1992, 1141–1142. Sato, S.; Matsunaga, K.; Horn, E.; Furukawa, N.; Nabeshima, T. J. Am. Chem. Soc. 2006, 128, 6778–6779. Sato, S.; Ameta, H.; Horn, E.; Takahashi, O.; Furukawa, N. J. Am. Chem. Soc. 1997, 119, 12374–12375. (a) Kumar, B.; Strasser, C. E.; King, B. T. J. Org. Chem. 2012, 77, 311–316; (b) El-Kurdi, S.; Seppelt, K. Chem. Eur. J. 2011, 17, 3956–3962. (a) Kleinhenz, S.; Schubert, M.; Seppelt, K. Chem. Ber. 1997, 130, 903–906; (b) Olmstead, M. M.; Power, P. P.; Shoner, S. C. Organometallics 1988, 7, 1380–1385; (c) Bartlett, R. A.; Power, P. P.; Shoner, S. C. J. Am. Chem. Soc. 1988, 110, 1966–1968. Fischer, R.; Görls, H.; Suxdorf, R.; Westerhausen, M. Organometallics 2019, 38, 498–511. Kinsella, E.; Smith, V. B.; Massey, A. G. J. Organomet. Chem. 1972, 34, 181–184. Usón, R.; Laguna, A.; Vicente, J. J. Chem. Soc. Chem. Commun. 1976, 353–354. Romão, C. C.; Kühn, F. E.; Herrmann, W. A. Chem. Rev. 1997, 97, 3197–3246. Herrmann, W. A.; Kuchler, J. G.; Felixberger, J. K.; Herdtweck, E.; Wagner, W. Angew. Chem. Int. Ed. 1988, 27, 394–396. Herrmann, W. A.; Ladwig, M.; Kiprof, P.; Riede, J. J. Organomet. Chem. 1989, 371, C13–C17. de Meric de Bellefon, C.; Herrmann, W. A.; Kiprof, P.; Whitaker, C. R, Organometallics 1992, 11, 1072–1081. Alves, A. S.; Moore, D. S.; Andersen, R. A.; Wilkinson, G. Polyhedron 1982, 1, 83–87. Herrmann, W. A.; Eder, S. J.; Kiprof, P.; Rypdal, K.; Watzlowik, P. Angew. Chem. Int. Ed. 1990, 29, 1445–1448. Tooze, R. P.; Stavropoulos, P.; Motevalli, M.; Hursthouse, M. B.; Wilkinson, G. J. Chem. Soc. Chem. Commun. 1985, 1139–1140. (a) Stravropoulos, P.; Edwards, P. G.; Behling, T.; Wilkinson, G.; Motevalli, M.; Hursthouse, M. B. J. Chem. Soc. Dalton. Trans. 1987, 169–175; (b) Longley, C. J.; Savage, P. D.; Wilkinson, G.; Hussain, B.; Hursthouse, M. B. Polyhedron 1988, 7, 1079–1088. Rankin, D. W. H.; Robertson, H. E.; Danopoulos, A. A.; Lyne, P. D.; Mingos, D. M. P.; Wilkinson, G. J. Chem. Soc. Dalton. Trans. 1994, 1563–1569. Muñiz, K. Chem. Soc. Rev. 2004, 33, 166–174. Hohenberger, J.; Ray, K.; Meyer, K. Nature Commun. 2012, 3. Article number 720. Poverenov, E.; Efremenko, I.; Frenkel, A. I.; Ben-David, Y.; Shimon, L. J. W.; Leitus, G.; Konstantinovski, L.; Martin, J. M. L.; Milstein, D. Nature 2008, 455, 1093–1096. (a) Schaub, S.; Miska, A.; Becker, J.; Zahn, S.; Mollenhauer, D.; Sakshath, S.; Schünemann, V.; Schindler, S. Angew. Chem. Int. Ed. 2018, 57, 5355–5358; (b) Schaub, S.; Miska, A.; Becker, J.; Zahn, S.; Mollenhauer, D.; Sakshath, S.; Schünemann, V.; Schindler, S. Angew. Chem. Int. Ed. 2019, 58, 5482. de Oliveira, F. T.; Chanda, A.; Banerjee, D.; Shan, X.; Mondal, S.; Que, L., Jr.; Bominaar, E. L.; Münck, E.; Collins, T. J. Science 2007, 315, 835–838. (a) Thorshaug, K.; Fjeldahl, I.; Rømming, C.; Tilset, M. Dalton Trans. 2003, 4051–4056; (b) Pellarin, K. R.; McCready, M. S.; Puddephatt, R. J. Dalton Trans. 2013, 42, 10444–10453. Meyer, S.; Klawitter, I.; Demeshko, S.; Bill, E.; Meyer, F. Angew. Chem. Int. Ed. 2013, 52, 901–905. Danopoulos, A. A.; Wilkinson, G.; Sweet, T. K. N.; Hursthouse, M. B. J. Chem. Soc. Dalton. Trans. 1994, 1037–1049. Danopoulos, A. A.; Wilkinson, G.; Sweet, T.; Hursthouse, M. B. J. Chem. Soc. Chem. Commun. 1993, 495–496. Spandl, J.; Supeł, J.; Drews, T.; Seppelt, K. Z. Anorg. Allg. Chem. 2006, 632, 2222–2225. Hursthouse, M. B.; Motevalli, M.; Sullivan, A. C.; Wilkinson, G. J. Chem. Soc. Chem. Commun. 1986, 1398–1399. Anneser, M. R.; Elpitiya, G. R.; Townsend, J.; Johnson, E. J.; Powers, X. B.; DeJesus, J. F.; Vogiatzis, K. D.; Jenkins, D. M. Angew. Chem. Int. Ed. 2019, 58, 8115–8118. Park, J. Y.; Kim, Y.; Bae, D. Y.; Rhee, Y. H.; Park, J. Organometallics 2017, 36, 3471–3476. Jacobs, B. P.; Wolczanski, P. T.; Jiang, Q.; Cundari, T. R.; MacMillan, S. N. J. Am. Chem. Soc. 2017, 139, 12145–12148. Campbell, M. M.; Johnson, G. Chem. Rev. 1978, 78, 65–79. Suzuki, H.; Nakaya, C.; Matano, Y.; Ogawa, T. Chem. Lett. 1991, 20, 105–108. Leeladee, P.; Jameson, G. N. L.; Siegler, M. A.; Kumar, D.; de Visser, S. P.; Goldberg, D. P. Inorg. Chem. 2013, 52, 4668–4682. Simkhovich, L.; Gross, Z. Tetrahedron Lett. 2001, 42, 8089–8092. Laplaza, C. E.; Cummins, C. C. Science 1995, 268, 861–863. Solari, E.; Da Silva, C.; Iacono, B.; Hesschenbrouck, J.; Rizzoli, C.; Scopelliti, R.; Floriani, C. Angew. Chem. Int. Ed. 2001, 40, 3907–3909. Caulton, K. G.; Chisholm, M. H.; Doherty, S.; Folting, K. Organometallics 1995, 14, 2585–2588. (a) Berry, J. F.; Bill, E.; Bothe, E.; DeBeer George, S.; Mienert, B.; Neese, F.; Wieghardt, K. Science 2006, 312, 1937–1941; (b) Nakamoto, K. Coord. Chem. Rev. 2002, 226, 153–165; (c) Meyer, K.; Bill, E.; Mienert, B.; Weyhermüller, T.; Wieghardt, K. J. Am. Chem. Soc. 1999, 121, 4859–4876; (d) Grapperhaus, C. A.; Mienert, B.; Bill, E.; Weyhermüller, T.; Wieghardt, K. Inorg. Chem. 2000, 39, 5306–5317. Betley, T. A.; Peters, J. C. J. Am. Chem. Soc. 2004, 126, 6252–6254. Bart, S. C.; Heinemann, F. W.; Anthon, C.; Hauser, C.; Meyer, K. Inorg. Chem. 2009, 48, 9419–9426. (a) Shapley, P. A.; Wepsiec, J. P. Organometallics 1986, 5, 1515–1517; (b) Shapley, P. A.; Kim, H. S.; Wilson, S. R. Organometallics 1988, 7, 928–933. Shapley, P. A.; Own, Z.-Y.; Huffman, J. C. Organometallics 1986, 5, 1269–1271. Shapley, P. A.; Schwab, J. J.; Wilson, S. R. J. Coord. Chem. 1994, 32, 213–232. Tonzetich, Z. J.; Lam, Y. C.; Müller, P.; Schrock, R. R. Organometallics 2007, 26, 475–477. de Frémont, P.; Mario, N.; Nolan, S. P. Coord. Chem. Rev. 2009, 253, 862–892. Schrock, R. R. J. Am. Chem. Soc. 1974, 96, 6796–6797. Guggenberger, L. J.; Schrock, R. R. J. Am. Chem. Soc. 1975, 97, 2935. (a) Andersen, R. A.; Chisholm, M. H.; Gibson, J. F.; Reichert, W. W.; Rothwell, I. P.; Wilkinson, G. Inorg. Chem. 1981, 20, 3934–3936; (b) Morton, L. A.; Zhang, X.-H.; Wang, R.; Lin, Z.; Wu, Y.-D.; Xue, Z.-L. J. Am. Chem. Soc. 2004, 126, 10208–10209. Savage, P. D.; Wilkinson, G.; Motevalli, M.; Hursthouse, M. B. Polyhedron 1987, 6, 1599–1601. Cheng, J.; Wang, L.; Wang, P.; Deng, L. Chem. Rev. 2018, 118, 9930–9987. Zhang, L.; Liu, Y.; Deng, L. J. Am. Chem. Soc. 2014, 136, 15525–15528. Hoppe, R. Rec. Trav. Chim. 1956, 75, 569–576. Martinez, G. E.; Ocampo, C.; Park, Y. J.; Fout, A. R. J. Am. Chem. Soc. 2016, 138, 4290–4293. (a) Vogel, C.; Heinemann, F. W.; Sutter, J.; Anthon, C.; Meyer, K. Angew. Chem. Int. Ed. 2008, 47, 2681–2684; (b) Scepaniak, J. J.; Vogel, C. S.; Khusniyarov, M. M.; Heinemann, F. W.; Meyer, K.; Smith, J. M. Science 2011, 331, 1049–1052.

134 130. 131. 132. 133. 134. 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. 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.

Very High Oxidation States in Organometallic Chemistry Herrmann, W. A.; Öfele, K.; Elison, M.; Kühn, F. E.; Roesky, P. W. J. Organomet. Chem. 1994, 480, C7–C9. Sundermeyer, J.; Weber, K.; Peters, K.; von Schnering, H. G. Organometallics 1994, 13, 2560–2562. Ellwanger, M. A.; Steinhauer, S.; Golz, P.; Braun, T.; Riedel, S. Angew. Chem. Int. Ed. 2018, 57, 7210–7214. Arduengo, A. J., III; Davidson, F.; Dias, H. V. R.; Goerlich, J. R.; Khasnis, D.; Marshall, W. J.; Prakasha, T. K. J. Am. Chem. Soc. 1997, 119, 12742–12749. Arduengo, A. J., III; Davidson, F.; Krafczyk, R.; Marshall, W. J.; Schmutzler, R. Monatsh. Chem. 2000, 131, 251–265. Connelly, N. G.; Geiger, W. E. Chem. Rev. 1996, 96, 877–910. Sharp, P. R.; Bard, A. J. Inorg. Chem. 1983, 22, 2689–2693. Malischewski, M.; Adelhardt, M.; Sutter, J.; Meyer, K.; Seppelt, K. Science 2016, 353, 678–682. (a) Lupinetti, A. J.; Frenking, G.; Strauss, S. H. Angew. Chem. Int. Ed. 1998, 37, 2113–2116; (b) Lupinetti, A. J.; Strauss, S. H.; Frenking, G. Progr. Inorg. Chem. 2001, 49, 1–112. Malischewski, M.; Seppelt, K.; Sutter, J.; Munz, D.; Meyer, K. Angew. Chem. Int. Ed. 2018, 57, 14597–14601. (a) Kölle, U.; Khouzami, F. Angew. Chem. Int. Ed. 1980, 19, 640–641; (b) Kölle, U.; Khouzami, F.; Lueken, H. Chem. Ber. 1982, 115, 1178–1196. Malischewski, M. Dissertation; Freie Universität Berlin, 2016. Struchkov, Y. T.; Antipin, M. Y.; Lyssenko, K. A.; Gusev, O. V.; Peganova, T. A.; Ustynyuk, N. A. J. Organomet. Chem. 1997, 536-537, 281–284. Herrmann, W. A.; Serrano, R.; Bock, H. Angew. Chem. Int. Ed. 1984, 23, 383–385. Dudeney, N.; Kirchner, O. N.; Green, J. C.; Maitlis, P. M. J. Chem. Soc. Dalton. Trans. 1984, 1877–1882. Jiménez-Tenorio, M.; Puerta, M. C.; Valerga, P. Organometallics 1994, 13, 3330–3337. Malischewski, M.; Seppelt, K.; Sutter, J.; Heinemann, F. W.; Dittrich, B.; Meyer, K. Angew. Chem. Int. Ed. 2017, 56, 13372–13376. Aneetha, H.; Tenorio, M. J.; Puerta, M. C.; Valerga, P. J. Organomet. Chem. 2002, 663, 151–157. Herrmann, W. A.; Okuda, J. Angew. Chem. Int. Ed. 1986, 25, 1092–1093. Herrmann, W. A.; Theiler, H. G.; Kiprof, P.; Tremmel, J.; Blom, R. J. Organomet. Chem. 1990, 395, 69–84. Gross, C. L.; Wilson, S. R.; Girolami, G. S. J. Am. Chem. Soc. 1994, 116, 10294–10295. Gilbert, T. M.; Bergman, R. G. Organometallics 1983, 2, 1458–1460. Isobe, K.; Bailey, P. M.; Maitlis, P. M. J. Chem. Soc. Chem. Commun. 1981, 808–809. (a) Crabtree, R. H. Acc. Chem. Res. 1990, 23, 95–101; (b) Hamilton, D. G.; Crabtree, R. H. J. Am. Chem. Soc. 1988, 110, 4126–4133. Brookhart, M.; Grant, B. E.; Lenges, C. P.; Prosenc, M. H.; White, P. S. Angew. Chem. Int. Ed. 2000, 39, 1676–1679. Ingleson, M.; Fan, H.; Pink, M.; Tomaszewski, J.; Caulton, K. G. J. Am. Chem. Soc. 2006, 128, 1804–1805. Fernandez, M. J.; Maitlis, P. M. J. Chem. Soc. Chem. Commun. 1982, 310–311. Fernandez, M. J.; Maitlis, P. M. Organometallics 1983, 2, 164–165. Chen, W.; Shimada, S.; Tanaka, M.; Kobayashi, Y.; Saigo, K. J. Am. Chem. Soc. 2004, 126, 8072–8073. Shimada, S.; Rao, M. L. N.; Tanaka, M. Organometallics 1999, 18, 291–293. (a) Shimada, S.; Tanaka, M.; Shiro, M. Angew. Chem. Int. Ed. 1996, 35, 1856–1858; (b) Shimada, S.; Tanaka, M.; Honda, K. J. Am. Chem. Soc. 1995, 117, 8289–8290. Chen, W.; Shimada, S.; Tanaka, M. Science 2002, 295, 308–310. (a) Crabtree, R. H. Science 2002, 295, 288–289; (b) Aullón, G.; Lledós, A.; Alvarez, S. Angew. Chem. Int. Ed. 2002, 41, 1956–1959; (c) Sherer, E. C.; Kinsingen, C. R.; Kormos, B. L.; Thompson, J. D.; Cramer, C. J. Angew. Chem. Int. Ed. 2002, 41, 1953–1956; (d) Nikonov, G. I. Angew. Chem. Int. Ed. 2003, 42, 1335–1337. Menjón, B.; Martínez-Salvador, S.; Gómez-Saso, M. A.; Forniés, J.; Falvello, L. R.; Martín, A.; Tsipis, A. Chem. Eur. J. 2009, 15, 6371–6382. Ramsden, C. A. Arkivoc 2014, 2014, 109–126. Tramšek, M.; Žemva, B. Acta Chim. Slov. 2006, 53, 105–116. Pérez-Bitrián, A.; Baya, M.l.; Casas, J. M.; Martín, A.; Menjón, B.; Orduna, J. Angew. Chem. Int. Ed. 2018, 57, 6517–6521. D’Accriscio, F.; Borja, P.; Saffon-Merceron, N.; Fustier-Boutignon, M.; Mézailles, N.; Nebra, N. Angew. Chem. Int. Ed. 2017, 56, 12898–12902. (a) Planas, O.; Wang, F.; Leutzsch, M.; Cornella, J. Science 2020, 367, 313–317; (b) Planas, O.; Wang, F.; Leutzsch, M.; Cornella, J. Science 2020, 367, eabb2416. Lal, G. S.; Pez, G. P.; Syvret, R. G. Chem. Rev. 1996, 96, 1737–1756. (a) Rozatian, N.; Ashworth, I. W.; Sandford, G.; Hodgson, D. R. W. Chem. Sci. 2018, 9, 8692–8702; (b) Gilicinski, A. G.; Pez, G. P.; Syvret, R. G.; Lal, G. S. J. Fluorine Chem. 1992, 59, 157–162. Meucci, E. A.; Ariafard, A.; Canty, A. J.; Kampf, J. W.; Sanford, M. S. J. Am. Chem. Soc. 2019, 141, 13261–13267. Engesser, T. A.; Lichtenthaler, M. R.; Schleep, M.; Krossing, I. Chem. Soc. Rev. 2016, 45, 789–899. Boeré, R. T.; Kacprzak, S.; Keßler, M.; Knapp, C.; Riebau, R.; Riedel, S.; Roemmele, T. L.; Rühle, M.; Scherer, H.; Weber, S. Angew. Chem. Int. Ed. 2011, 50, 549–552. Zemva, B. C. R. Acad. Sci. Paris 1998, 1, 151–156. Wizansky, A. R.; Rauch, P. E.; Disalvo, F. J. J. Solid State Chem. 1989, 81, 203–207. Poleschner, H.; Seppelt, K. Angew. Chem. Int. Ed. 2013, 52, 12838–12842. Steckhan, E. Top. Curr. Chem. 1987, 142, 1–69. Reed, C. A.; Kim, K.-C.; Bolskar, R. D.; Mueller, L. J. Science 2000, 289, 101–104. Schorpp, M.; Heizmann, T.; Schmucker, M.; Rein, S.; Weber, S.; Krossing, I. Angew. Chem. Int. Ed. 2020, 59, 9453–9459. (a) Malinowski, P. J.; Himmel, D.; Krossing, I. Angew. Chem. Int. Ed. 2016, 55, 9259–9261; (b) Malinowski, P. J.; Himmel, D.; Krossing, I. Angew. Chem. Int. Ed. 2016, 55, 9262–9266. (a) Aynsley, E.; Peacock, R. D.; Robinson, P. L. Chem. Ind. (London) 1951, 1117; (b) Moore, J. W.; Baird, H. W.; Miller, H. B. J. Am. Chem. Soc. 1968, 90, 1358–1359. Byler, D. M.; Shriver, D. F. Inorg. Chem. 1974, 13, 2697–2705. Ferron, B.; Jacquesy, J.-C.; Jouannetaud, M.-P.; Karam, O.; Coustard, J.-M. Tetrahedron Lett. 1993, 34, 2949–2952. Seppelt, K. Chem. Rev. 2015, 115, 1296–1306. Macgregor, S. A.; Moock, K. H. Inorg. Chem. 1998, 37, 3284–3292. (a) Bartlett, N. Proc. Chem. Soc. 1962, 218; (b) Bartlett, N.; Lohmann, D. H. Proc. Chem. Soc. 1962, 115–116. Dinnocenzo, J. P.; Banach, T. E. J. Am. Chem. Soc. 1986, 108, 6063–6065.

1.06

Characterization Methods for Paramagnetic Organometallic Complexes

Aleksa Radovic, Shilpa Bhatia, and Michael L Neidig, Department of Chemistry, University of Rochester, Rochester, NY, United States © 2022 Elsevier Ltd. All rights reserved.

1.06.1 1.06.2 1.06.2.1 1.06.2.2 1.06.2.3 1.06.2.3.1 1.06.2.3.2 1.06.2.4 1.06.3 1.06.3.1 1.06.3.2 1.06.3.3 1.06.4 1.06.4.1 1.06.4.2 1.06.4.3 1.06.5 1.06.5.1 1.06.5.2 1.06.5.3 1.06.6 References

1.06.1

Introduction Electron paramagnetic resonance (EPR) spectroscopy Introduction Theory Continuous wave (CW) EPR spectroscopy X-band EPR spectroscopy High field EPR Pulsed EPR spectroscopy Magnetic circular dichroism (MCD) spectroscopy Introduction Theory Applications X-ray absorption spectroscopy Introduction Theory Applications Nuclear magnetic resonance Introduction Theory Applications Conclusion

135 136 136 136 139 139 145 147 148 148 148 150 155 155 156 158 165 165 165 169 172 172

Introduction

Electronic structure and bonding are central to the physical properties and reactivities of organometallic compounds.1,2 For organometallic compounds of precious metals as well as base metal organometallic compounds with strong field ligands (e.g., CO ligand) that have historically dominated organometallic chemistry, the ground states are singlets, and the diamagnetism has facilitated extensive insight into these key properties using traditional characterization methods, such as NMR spectroscopy.3,4 More recently, base metal organometallics have become a major research focus, motivated by the goal of developing more sustainable and economical systems as well as accessing new reaction manifolds.5 While these systems have proven highly successful in terms of the catalytic efficiency and are very advantageous due to the abundance and cost efficiency of the involved metals, a fundamental understanding of the contributions of electronic structure and bonding to their physical properties and reactivities remains underdeveloped compared to their precious metal analogs. Central to this challenge is the prevalence of paramagnetic complexes and intermediates in base metal organometallic chemistry and the associated characterization challenges resulting from the presence of unpaired electrons in the ground states of such systems.6,7 Numerous characterization methods can provide critical insight into paramagnetic organometallic systems. For example, common techniques such as infrared spectroscopy, electronic absorption spectroscopy and magnetic measurements can provide important information on the oxidation and spin state of paramagnetic organometallics.8–11 Magnetic measurements are also particularly useful in studies of multinuclear systems for which ferromagnetic or anti-ferromagnetic coupling may exist.12,13 These characterization methods can be further supplemented by the use of resonance Raman spectroscopy or Mössbauer spectroscopy (most commonly used for iron complexes) to obtain further insight into bonding and electronic structure in organometallic complexes.14–17 While the aforementioned techniques are useful for the characterization of paramagnetic organometallic complexes, this chapter is focused on four specific spectroscopic methods that have proven especially effective in evaluating electronic structure and bonding in these complexes. Specifically, these are electron paramagnetic resonance spectroscopy (EPR), magnetic circular dichroism spectroscopy (MCD), X-ray absorption spectroscopy (XAS) and nuclear magnetic resonance spectroscopy (NMR). EPR is a powerful tool for analysis of systems which contain unpaired electrons, like paramagnetic organometallic complexes.18 In this technique, transitions between unpaired electron spin energy levels are probed under the influence of external magnetic field. One of the most common uses of EPR is determination of the ground spin state from which oxidation state of the metal center can be inferred. As unpaired electron spin density is directly probed, EPR is very useful for determination of electronic structure and geometry of the metal sites in organometallic complexes. The high sensitivity of EPR makes it ideal for characterization of in situ formed paramagnetic organometallic intermediates.19 Also, it has proven to be useful in analysis of systems with redox active

Comprehensive Organometallic Chemistry IV

https://doi.org/10.1016/B978-0-12-820206-7.00059-7

135

136

Characterization Methods for Paramagnetic Organometallic Complexes

ligands, as it can be used to determine is unpaired electron density located on metal center or on ligand. The shape of the peak in EPR is influenced by the interactions of an unpaired electron with its environment. The nature of this interaction can yield important information, such as rates of chemical reactions.20 MCD spectroscopy is another technique that can provide tremendous insight into the electronic structure in paramagnetic organometallic systems.21–23 It is based on the differential absorption of left and right circularly polarized light, induced in a sample by a strong magnetic field oriented parallel to the direction of light propagation, and it is applicable to any S > 0 organometallic complex. MCD measurements can detect transitions which are often too weak to be observed in conventional optical absorption spectra (d-d and f-f transitions). Additionally, the presence of positive and negative transitions can be used to distinguish between overlapping transitions which are often broad in absorption spectroscopy.24,25 MCD can also provide insight into the structure of organometallic compounds, including insight into coordination number, geometry and ground state spin Hamiltonian parameters including zero-field splitting (ZFS), from which the oxidation state of a metal within a complex can be inferred.26 XAS spectroscopy is based on the absorption of X-ray radiation which can lead to excitation of core electrons to the lowest unoccupied orbitals or ionization.27 This technique is element-specific which means that metal centers in organometallic complexes can be directly probed. Complementary information can be obtained from the ligand elements. XAS is mostly used for determination of oxidation state of metal centers, geometry, electronic structure and covalency. It can also be used for studying local structure through the number, type and distance of neighboring atoms.28–30 The analysis of the XANES (X-ray absorption near edge structure) region can provide valuable information about oxidation states and coordination environments, which can be particularly useful when studying systems with redox active ligands.31 XAS is discussed in this chapter to facilitate proper analysis of paramagnetic systems, as understanding of some details is crucial for accurate conclusions. NMR spectroscopy of paramagnetic compounds can also provide significant amount of useful information.4 The interaction of unpaired electrons with NMR active nuclei results in broadening and shifts of resonances beyond that traditionally observed for diamagnetic compounds. One of the most-used aspects of paramagnetic NMR is determination of effective magnetic moments for paramagnetic compounds, from which the spin of the ground state can be determined.32,33 This technique has been utilized in structure elucidation and characterization of organometallic compounds, including intermediates formed in chemical transformations.34,35 In this chapter, each of these spectroscopic techniques will be discussed in detail with respect to their use in characterizing paramagnetic organometallic complexes. Fundamentals of each spectroscopic method are presented to provide the necessary background and theory central to each technique. This is followed by discussion of the key applications of each method for paramagnetic organometallic complexes, using specific examples from the literature to highlight the insight into electronic structure, bonding, and reactivity.

1.06.2

Electron paramagnetic resonance (EPR) spectroscopy

1.06.2.1

Introduction

Electron paramagnetic resonance (also known as electron spin resonance, ESR) spectroscopy is one of the most accessible and widely utilized techniques for studying paramagnetic organometallic complexes. This technique is based on the absorption of electromagnetic radiation by a paramagnetic sample placed in an external applied magnetic field. The magnetic field removes the degeneracy of unpaired electron spin energy levels via the Zeeman effect and transitions between these Zeeman split levels are probed. EPR is limited to studies of paramagnetic species including organic radicals as well as d- and f-element compounds with unpaired electrons. In organometallic chemistry, EPR is widely employed for the determination of the ground state spin, from which the metal oxidation state can often be inferred. Furthermore, it can provide valuable insight into the electronic structure and bonding of organometallic compounds, including metal-ligand covalency. Lastly, this method can be utilized for mechanistic studies, including the identification of in situ formed reaction intermediates such as those formed in one-electron pathways of organometallic complexes.

1.06.2.2

Theory

The development of this technique followed from the Stern-Gerlach experiment, which showed that electrons possess intrinsic angular momentum called spin, S.36–38 Since the electron is a charged particle, the spin gives rise to a magnetic moment (m): m ¼ − ge mB S

(1)

where mB is a Bohr magneton and ge is an electron g factor (for free electron ge  2.0023). The spin of the electron is equal to 1/2 in reduced Planck’s constant units (ħ), and it can have two different states with respect to the projection of the spin to the z-axis in the Cartesian system. These two states are called the a state, where Sz ¼ +1/2, and the b state, where Sz ¼ − 1/2 (in ħ units). In the absence of an external magnetic field these two states are degenerate. Application of an external magnetic field results in an interaction between this field and magnetic moment of the electron, represented by the electron Zeeman interaction Hamiltonian:

Characterization Methods for Paramagnetic Organometallic Complexes

137

b Ze ¼ − mB ¼ ge mB Bb H S

(2)

E ¼ ge mB BMs

(3)

where B represents the vector of the applied magnetic field. Since it can be chosen that the magnetic field vector is aligned along the z-axis, and using the relation Sz ¼ Msħ, where Ms is magnetic spin quantum number (Ms ¼ S, S − 1, . . ., −S), the energy of the electron Zeeman interaction is:

From this equation, states with different Ms values have different energies, and in the case of a S ¼ 1/2 system the energy difference between these states is: DE ¼ ge mB B

(4)

From the resonance condition, the energy difference from Eq. (4) must be equal to the energy of applied electromagnetic radiation in order to observe a transition between the Zeeman split states (Fig. 1). The orientation of the magnetic field component of the electromagnetic radiation (B1) toward the applied static magnetic field (B) governs the selection rules for the transition. In the case of a longitudinal configuration, where B1 is parallel to B, the selection rule is DMs ¼ 0 (parallel mode EPR). By contrast, for the transverse configuration where B1 is perpendicular to B, the selection rule is DMs ¼  1 (perpendicular mode EPR). In metal complexes, the g factor is a property of the spin system, not a specific electron, and therefore differs from the free electron value. In this case the g factor is a 3  3 tensor that describes the orientation-dependent response of the energy levels to a magnetic field. Since g is a tensor, it can be represented with three components along the principal axes gx, gy, and gz (most often represented as g ¼ [gx gy gz]). In a metal complex, the angular momentum of the unpaired electron is affected by orbital angular momentum through the coupling of the spin and orbital angular momenta, which is known as spin-orbit coupling (SOC). Consequently, an effective spin Hamiltonian (which contains just the electron and nuclear spin operators) is used to give an

Fig. 1 Energy levels and EPR transitions for S ¼ 1/2, I ¼ 1/2 and IL ¼ 1/2 system.

138

Characterization Methods for Paramagnetic Organometallic Complexes

appropriate description of these systems. Since the magnitude of the SOC depends on the SOC constant of a given atom, which increases with atomic number Z, organic free radicals have g values (factors) close to the ge value due to the small SOC constants of C, N, O, etc. By contrast, organometallic compounds can have g values that significantly deviate from ge due to the relatively large SOC from the metal.39 In the case of a system where electronic distribution has cubic symmetry, SOC is isotropic meaning that all g components are equivalent (isotropic signal). By contrast, in axial symmetry gx ¼ gy 6¼ gz meaning that there are two different g values. These two g values are often referred as g‖ (¼gx ¼ gy) and g? (¼gz) (axial signal). In rhombic symmetry, all three g values are unique (rhombic signal). This anisotropy can be used to provide information about the symmetry of the electronic distribution in an investigated system. Note that while EPR is equally applicable to both solid samples (solid or frozen solution) and solution samples, for the latter only an averaged g value can be observed due to the rapid tumbling of molecules in solution.39,40 The unpaired electron can also interact with surrounding nuclei with net nuclear spin (I > 0). This interaction is called a hyperfine interaction and can be represented by the hyperfine interaction Hamiltonian36,41: b hf ¼ Ab H SbI

(5)

where A is the hyperfine coupling constant. As a result, the degeneracy with respect to the magnetic nuclear spin quantum number MI is removed and, since there are 2I + 1 possible MI values (MI ¼ I, I − 1, . . ., −I), the Zeeman transitions will be split to 2nI + 1 lines (n is number of equivalent nuclei) with a separation of A (Fig. 1).42 Note that an additional selection rule in this case is DMI ¼ 0. Since this splitting is small, it is expected that all levels with the same MS value will be equally populated and, hence, all lines will have equal intensity. The hyperfine interaction is a consequence of three interactions between the electron and nucleus magnetic moments: Fermi contact interaction, direct and indirect dipolar interaction. The Fermi contact and indirect dipolar interactions give rise to isotropic hyperfine coupling, while the direct dipolar coupling is orientation dependent and leads to anisotropic hyperfine coupling. Note that while the interaction of the unpaired electron with metal nuclei is called the metal hyperfine interaction, interaction with ligand nuclei is alternatively termed the ligand hyperfine or superhyperfine interaction. The nature of these interactions is the same, but the superhyperfine interaction is typically an order of magnitude smaller than the hyperfine interaction. Even though weak, the metal hyperfine interactions can in some cases be observed in the most commonly used X-band EPR experiments (X-band refers to the frequency of microwaves used in experiment, as shown later in this chapter), while the weaker ligand hyperfine interactions are not readily observed (except in some favorable situations as shown later in this chapter). Instead, these interactions are generally studied by pulsed EPR methods, which have higher resolution than continuous wave X-band EPR and thus enable direct observation of these weak interactions. In systems with more than one unpaired electron (S > 1/2) there is an additional effect called zero-field splitting (ZFS).36 ZFS removes the degeneracy of the ground state MS levels, even in the absence of an external magnetic field. In the case of high symmetry (octahedral or higher) organometallic complexes, SOC is isotropic and there is no ZFS. Lowering of the symmetry results in anisotropic SOC, which leads to ZFS. This interaction can be represented with the ZFS Hamiltonian:     1 b ZFS ¼ b H (6) SDb S ¼ D S2z − SðS + 1Þ + E S2x − S2y 3 2Dz − ðDx + Dy Þ where D is the zero-field interaction tensor, D is the axial ZFS parameter (D ¼ ) and E is the rhombic ZFS parameter 2 Dx − Dy (E ¼ 2 ). For cubic symmetry, D ¼ E ¼ 0 and there is no ZFS. For axial symmetry D 6¼ E ¼ 0, and for rhombic symmetry D 6¼ E 6¼ 0. The relation between D and E determines the overall influence on g and is defined along an axis such that43: jEj 1 ¼l jDj 3

(7)

While for axial symmetry systems l ¼ 0, increasing rhombicity (e.g., lowering symmetry) increases the value of l by up to 1/3 for purely rhombic systems. As a representative example, Fig. 2 shows the splitting of ground-state Ms levels for different symmetries for a S ¼ 1 system. Lowering symmetry from cubic to axial leads to the splitting of |Ms | ¼ 1 and Ms ¼ 0 states by D. Additional lowering of the symmetry to rhombic leads to removal of the degeneracy of the Ms ¼ 1 and Ms ¼ −1 levels with a splitting of 2E. Besides increasing the complexity of EPR spectra, ZFS parameters can provide important information about structure and bonding within the organometallic complex of interest.44 Lastly, systems having nuclei with I  1 can have an additional interaction, which is a consequence of the non-spherical charge distribution in these nuclei, resulting in a non-zero nuclear quadrupole moment. This quadrupole moment may interact with an electric field gradient (from nearby electrons and nuclei), termed the nuclear quadrupole interaction.36,39 The Hamiltonian for this interaction is characterized by the nuclear quadrupole interaction tensor. While this quadrupolar interaction is weak and higher resolution EPR techniques are necessary to observe its effect on the resulting spectra, it can provide valuable information about electronic structure and bonding in organometallic complexes. Systems with more than one unpaired electron have so far been treated as strongly interacting electrons, and the total spin of the system is used to describe all interactions. In multinuclear organometallic complexes it is possible that unpaired electrons on different paramagnetic centers are weakly coupled, interacting through exchange coupling and dipole-dipole coupling.36 Exchange coupling becomes important when there is significant overlap of orbitals which contain unpaired electrons at different metal centers. In solid samples, exchange coupling becomes significant if unpaired electrons on different centers are 15 A˚ or closer apart or highly delocalized.39 The interacting spins can be aligned, which corresponds to the ferromagnetic coupling (for two interacting

Characterization Methods for Paramagnetic Organometallic Complexes

139

Fig. 2 ZFS of energy levels for S ¼ 1 system for cubic, axial and rhombic symmetries.

spins, S ¼ S1 + S2) or anti-aligned which corresponds to the antiferromagnetic coupling (for two interacting spins, S ¼ | S1 − S2|). The strength of this interaction is given by an exchange constant J, where the sign of this constant is positive if spins are ferromagnetically coupled and negative if spins are antiferromagnetically coupled. The dipole-dipole coupling between individual spins is analogous to dipole-dipole coupling between electron and nuclear spins. This interaction contains information about the distance between interacting spins. For organometallic compounds, the most important exchange coupling is through the bridging ligands between metal centers. Relaxation processes also play an important role in EPR spectroscopy. In the simplest case (e.g., S ¼ 1/2 system with only electron Zeeman interaction), the population of states is governed by the Boltzmann distribution. Since thermal energy (at room temperature) is much larger than the Zeeman splitting, both MS states are almost equally populated, with only a small excess of spins in the lower energy state. After the absorption of electromagnetic radiation of the appropriate energy, this excess of spins is moved from the ground to the excited state (“excitation”) and the system evolves over time to return to the ground state (“relaxation”). The system cannot absorb the same electromagnetic radiation until excited spins are relaxed back to the ground state. In order to achieve larger population difference in split states, EPR spectra of organometallic complexes are often obtained at low temperatures (e.g., liquid He temperature). This approach reduces broadening of bands in the resulting spectra. Additionally, in some cases it is not possible to observe bands corresponding to S > 1/2 states at higher temperatures (e.g., like in the case of some high-spin iron(III) complexes).

1.06.2.3

Continuous wave (CW) EPR spectroscopy

In order to observe a transition in EPR, the energy of the applied electromagnetic radiation must satisfy the resonance condition (Eq. 4). In most cases this is achieved by varying the magnetic field at a constant microwave frequency. Since absorption is weak due to the small population difference of Zeeman split sublevels, the signal must be detected using a resonator. The detection mechanism is responsible for the line shape of the EPR spectrum, which is shown as the first derivative of the absorption spectrum. This line shape also makes it easier to detect and quantify partially resolved splitting in EPR spectra.45 Most EPR spectra are recorded at microwave frequencies 9.5 GHz (X-band).42 This is the most commonly used frequency because the magnetic field required to achieve resonance at this frequency can be achieved using electromagnet and microwave components that are relatively cheap while at the same time providing high sensitivity. If features in the EPR spectra are unresolved, or ZFS is too large for X-band microwaves (0.3 cm−1), higher field EPR can be used. All commonly used frequencies are shown in Table 1. Higher frequencies require larger magnetic fields and smaller resonators (less sample). W-band (and higher frequency) instruments require superconducting magnets which significantly increases the cost and complexity of the spectrometer.45 Ultimately, the choice of appropriate frequency for the EPR experiment depends on the properties of the investigated systems. The samples for continuous wave EPR spectroscopy can be in solution or in solid state, where the latter also includes frozen solutions. As mentioned earlier, the spectra of organometallic complexes are most often taken at low temperatures. This enables the analysis of chemical transformations by freezing reaction mixtures at specific timepoints, and it can also be used for studying species that are unstable at higher temperatures.

1.06.2.3.1

X-band EPR spectroscopy

As mentioned above, X-band CW EPR spectroscopy in the perpendicular mode configuration is the most widely used EPR technique in organometallic chemistry. For example, it can be used for the identification of compounds (in combination with other techniques), the determination of spin state, symmetry of molecules and ZFS parameters. In some cases, it can be used to determine

140

Characterization Methods for Paramagnetic Organometallic Complexes

Table 1

Commonly used frequencies and magnetic fields for g ¼ 2.0.

Band

Frequency (GHz)

Magnetic field (T)

L S X Q W High field

1.5 2–4 9–10 33–35 95 >95

0.05 0.07–0.14 0.32–0.36 1.20–1.30 3.4 >3.4

Adapted from Schweiger, A.; Jeschke, G. Principles of Pulse Electron Paramagnetic Resonance; Oxford University Press, 2001.

hyperfine and superhyperfine couplings, but often these interactions are too weak to be resolved by this technique. Perpendicular mode X-band EPR is limited to half-integer spin systems (Kramers doublets), because these systems have at least doubly degenerate ground states where perpendicular mode EPR transitions are allowed. In order to analyze integer spin (non-Kramers) systems, parallel mode X-band EPR can be used as it has different selection rule for EPR.42,45 However, even in parallel mode it would be only possible to see transitions for integer spin systems in which the rhombic ZFS parameter is non-zero and smaller than the energy of X-band microwaves (0.3 cm−1). This limitation significantly reduces the number of systems that can be analyzed using parallel mode EPR spectroscopy.46 This means that the absence of a signal in both perpendicular and parallel mode EPR spectra does not necessarily mean that the system is diamagnetic. Perpendicular mode X-band EPR can be used for characterizing the ground state spin of organometallic compounds with Kramers ground states, quantifying the amount of such species present in solution and following the in situ generation or consumption of species during reactions or catalysis.47–55 In addition, this method is especially useful regarding the determination of the ground state spin of an organometallic complex for which multiple different spin states are possible for a given d-electron count. For example, Neidig and co-workers used perpendicular mode EPR in order to characterize the ground-state spin of an unusual distorted square-planar tetramethyl iron(III) ferrate complex [MgCl(THF)5][FeMe4]
THF.56 Since iron(III) is a half-integer spin system (d5), multiple different ground spin states are possible: S ¼ 1/2 (low spin), 3/2 (intermediate spin) and 5/2 (high spin). As can be seen from Fig. 3, this iron complex has a characteristic S ¼ 3/2 axial EPR signal, with two signals at 1500 (g  4.20) and 3300 (g  2.00) Gauss.57–59 Generally, the resolution of spectra in X-band EPR is not sufficient to allow direct simulation of experimental data in order to extract ZFS parameters in S > 1/2 systems. In this example the spectrum was recorded at several different temperatures (Fig. 3) and the normalized intensity data were fit to a Curie law dependent Boltzmann distribution in order to extract ZFS parameters. This particular study also serves as an example of the use of perpendicular mode X-band EPR to follow in situ reactions of iron species. Furthermore X-band EPR of samples freeze-trapped at different times was also used to monitor the reduction of this iron(III) complex to a new S ¼ 1/2 species. Fig. 4 shows EPR spectra after different warming times, which resulted in a new, broad feature at g  2 with concomitant loss of the S ¼ 3/2 signal. This S ¼ 1/2 product was later confirmed by a combination of X-ray crystallography and EPR spectroscopy to be a [Fe8Me12]− cluster.60 A notable feature of EPR spectroscopy is that it can be used to determine the concentration of paramagnetic organometallic species through comparison between the integrated intensity (area under the peak of the absorption spectrum; double integration of the derivative spectrum) of an EPR signal and that of an external standard. For such studies, it is important that the spectra of the investigated species and the standard are taken under identical conditions. If the spectra are not taken under identical conditions the intensities can often be adjusted to take into account the differences in instrumental conditions. This strategy was used in the

Fig. 3 X-band EPR spectrum of [MgCl(THF)5][FeMe4]
THF complex in THF (10 K). Insert shows temperature dependence of signal intensity, which is fitted to the Curie law. Reprinted with permission from Al-Afyouni, M. H.; Fillman, K. L.; Brennessel, W. W.; Neidig, M. L. J. Am. Chem. Soc. 2014, 136, 15457–15460. Copyright 2014 American Chemical Society.

Characterization Methods for Paramagnetic Organometallic Complexes

141

Fig. 4 X-band EPR spectra (10 K) of (A) [MgCl(THF)5][FeMe4]
THF in THF at −80  C, (B–D) after warming to −40  C at different time points, (E) after warming to room temperature. Reprinted with permission from Al-Afyouni, M. H.; Fillman, K. L.; Brennessel, W. W.; Neidig, M. L. J. Am. Chem. Soc. 2014, 136, 15457–15460. Copyright 2014 American Chemical Society.

previous example, where the authors wanted to determine: (1) the amount of iron converted to [MgCl(THF)5][FeMe4]
THF upon reaction of FeCl3 with Grignard reagent and (2) the amount of S ¼ 1/2 species formed after the warming of [MgCl(THF)5][FeMe4]
THF.56 The EPR spectrum of the S ¼ 1/2 species was integrated using a 1 mM solution of CuSO4 standard with identical instrumentation parameters for both iron and standard samples. For the S ¼ 3/2 iron species, the integration was performed with the same standard following the established method of Aasa and Vänngård61,62 and corrected for the Boltzmann distribution in the ground state spin manifold using the experimentally determined zero-field spitting parameter for S ¼ 3/2 iron species. The Boltzmann correction was necessary as the standard and sample have different spin states. Based on the results the authors determined that 50  10% of iron was converted in the initial reaction to [MgCl(THF)5][FeMe4]
THF, while 95  10% of the final S ¼ 1/2 iron product was formed upon warming [MgCl(THF)5][FeMe4]
THF. X-band perpendicular mode EPR spectroscopy can be used for characterizing organometallic compounds containing any metal, as long as the complex has a half-integer total spin. One area that has received broad research interest is the EPR studies of Ni(I) and Ni(III) organometallic species, motivated by their proposed roles in numerous reactions catalyzed by nickel.63–67 In one representative example, Mirica and co-workers synthesized and characterized Ni(III) (d7) complexes that can undergo C-C or C-heteroatom

142

Characterization Methods for Paramagnetic Organometallic Complexes

Fig. 5 (A) EPR spectrum (77 K) of 1+ and 2+ (red) and corresponding simulations (blue), (B) EPR spectrum (77 K) of 4+ and 13C-4+ (red) and corresponding simulations (blue). Adapted with permission from Zheng, B.; Tang, F.; Luo, J.; Schultz, J. W.; Rath, N. P.; Mirica, L. M. J. Am. Chem. Soc. 2014, 136, 6499–6504. Copyright 2014 American Chemical Society.

bond forming reactions.68 X-band EPR spectroscopy was used for characterization of [(tBuN4)NiIII(PhF)X]+ (PhF ¼ p-fluorophenyl; X ¼ Br: 1+; Cl: 2+; CH3: 4+; 13CH3: 13C-4+). Complexes 1+ and 2+ adopt distorted octahedral geometries around the metal center. Magnetic susceptibility measurements indicated that these complexes contain one unpaired electron (S ¼ 1/2). The EPR spectra (frozen solution, 77 K) of complexes 1+ and 2+ (Fig. 5A) were characterized by a rhombic signal at the average g value of 2.160–2.162, which suggested that the unpaired electron is located on the metal center (as the g value significantly differs from the ge value). Also, a superhyperfine interaction with two axially bound nitrogens (I ¼ 1 for 14N) was observed in the splitting of the gz component, and for complex 1+ the additional superhyperfine interaction with Br (I ¼ 3/2 for both stable isotopes) can be observed in the gx and gy components. These data were most consistent with an unpaired electron localized mostly on the metal dz2 orbital for these distorted octahedral Ni(III) complexes. The complexes where the halide was replaced with a methyl group show similar features in the EPR spectrum (Fig. 5B) with a superhyperfine interaction from axial nitrogens along the z direction. Additionally, in the complex where a 13C labeled methyl group was used, additional superhyperfine interaction with 13C (I ¼ 1/2) is observed along x and y directions, through the broadening of the gx and gy signals (as the magnetic field in X-band EPR is not sufficient to cause complete splitting of the lines). Superhyperfine interaction along x and y axes is a consequence of the contribution of the 3dx2−y2 orbital to the singly occupied molecular orbital. Based on previous findings, it was concluded that the SOMO consists mostly of metal dz2 orbital with additional contribution from the metal 3dx2−y2 orbital. In another example, Hazari, Nova and coworkers used X-band EPR spectroscopy to study (Z5-Cp)NiI(IPr) (Cp ¼ cyclopentyldienyl, IPr ¼ 1,3-bis(2,6diisopropylphenyl)-1,3-dihydro-2H-imidazol-2-ylidene).69 The EPR spectrum shows a rhombic signal in accordance with previously studied Ni(I) complexes with S ¼ 1/2 ground state.70 It was also observed that similar features were present in the EPR spectra of both Ni(I) and Ni(III) organometallic complexes, indicating that EPR should ideally be used in combination with other techniques in order to unambiguously determine the oxidation state of Ni.

Characterization Methods for Paramagnetic Organometallic Complexes

143

Fig. 6 (A) Structure of studied Co complex, (B) X-band EPR-monitored photolysis of cobalt complex in 2-methyl-THF solution (10 K). Adapted with permission from Zolnhofer, E. M.; Käß, M.; Khusniyarov, M. M.; Heinemann, F. W.; Maron, L.; van Gastel, M.; Bill, E.; Meyer, K. J. Am. Chem. Soc. 2014, 136, 15072–15078. Copyright 2014 American Chemical Society.

X-band EPR is also useful for the characterization of unstable paramagnetic intermediates, which cannot be isolated and crystallized. In the work of Meyer and coworkers, EPR spectroscopy was used for studying the photolysis of a Co(II) azide complex, supported by an NHC containing multidentate ligand (Fig. 6A), leading to insertion of an N into the M–C bond.71 The resulting product was characterized by a low temperature EPR. The broad axial signal at g? ¼ 4.18 and g‖ ¼ 2.02 corresponds to a Co(II) complex in an S ¼ 3/2 ground state. Due to the broadness of the signal, the hyperfine splitting from the 59Co coupling could not be observed. It was proposed that this reaction proceeds via the formation of a Co(IV) nitrenoid intermediate, and it was possible to trap this transient species by irradiating the initial complex inside the EPR cryostat, followed by X-band EPR spectroscopy measurement at 10 K. Fig. 6B shows the in situ spectrum after photolysis. The EPR active intermediate is characterized by an isotropic signal at g ¼ 2.01 with 8-line due to hyperfine coupling to the 59Co nucleus (I ¼ 7/2). The observed signal agrees with the predicted spectrum for a low spin (S ¼ 1/2) Co(IV) complex in a trigonal ligand field. Superhyperfine coupling to the N nuclei, which would provide definite evidence of formed nitrenoid, could not be observed due to the broadness of the signal. Despite this complication, the observations from low temperature EPR measurements support the proposed formation of Co(IV) complex from the precursor Co(II) complex by loss of N2. EPR at a higher temperature (77 K) showed rapid decay, indicating that the formed intermediate is unstable at this slightly elevated temperature and showing the exceptional ability of EPR for characterization under conditions that could not be used for other spectroscopic techniques. The characterization of organometallic complexes by EPR has also proven useful in establishing the mechanisms for various transformations that involve the formation of paramagnetic organometallic intermediates.19,72–75 For example, Diao and coworkers used EPR for the characterization of intermediates in Ni catalyzed C–C cross-coupling reactions, focusing on the formation of ethane from (py-Mepyrr)Ni(CH3)(lut) (py-Mepyrr ¼ 3,5-dimethyl-2-(2-pyridyl)pyrrole, lut ¼ 2,4-lutidine).76 In this example it had been shown that employing either I2 or NBS (N-bromosuccinimide) as oxidants give highest yields for the formation of ethane. Treatment of this Ni complex with I2 gave a new species with an EPR spectrum (Fig. 7) containing a S ¼ 1/2 signal with observable superhyperfine splitting in the z direction This superhyperfine interaction gives a three-line pattern (1:1:1 splitting) that originates from the interaction of an unpaired electron with N nuclei (I ¼ 1) on the pyridine ligand in axial position. Additionally, the EPR spectrum of the mixture after standing at −20  C for 10 min contained no EPR signal, suggesting that this intermediate was unstable. Switching from I2 to NBS resulted in a slight change in the EPR spectrum with a shift in the g-values (presumably because of the difference between I and Br), but similar spectral features were observable. This suggested that the complexes formed with I2 and NBS have similar structures. DFT calculations led to the proposal of (py-Mepyrr)NiIII(X)(CH3)(lut) (X ¼ I, Br) as a possible structure of the spectroscopically observed intermediates. Overall, the reaction mechanism proposed from these studies involves the dissociation of 2,4-lutidine from the EPR active Ni(III) complex, resulting in the formation of the dinuclear Ni(III) complex which yields ethane upon reductive elimination. In studies of organometallic systems with potentially redox non-innocent ligands, the determination of the oxidation state of the metal center can be challenging, because the unpaired electron could lie on the redox non-innocent ligand rather than on the metal (as would be assumed using the formal oxidation state). It is worth mentioning again that the EPR experiment probes the spin of the whole system, not local spin. For example, this means that an Fe(III) complex is not active in perpendicular mode X-band EPR if it contains a radical anion ligand, as the total number of unpaired electrons is even. EPR spectroscopy can be useful for the characterization of systems with redox non-innocent ligands, as it can provide information about whether unpaired electrons are localized on the metal center or on the ligand.77–80 An example of the application of EPR spectroscopy in organometallic systems containing redox non-innocent ligands can be found in the work of Milstein, Neidig and coworkers that focused on

144

Characterization Methods for Paramagnetic Organometallic Complexes

Fig. 7 X-band EPR spectra of intermediates formed after reaction of (py-Mepyrr)Ni(CH3)(lut) with I2 (red) and NBS (black line) in toluene (10 K) and corresponding simulation (blue). Reprinted with permission from Xu, H.; Diccianni, J. B.; Katigbak, J.; Hu, C.; Zhang, Y.; Diao, T. J. Am. Chem. Soc. 2016, 138, 4779–4786. Copyright 2016 American Chemical Society.

a-iminopyridine based iron-PNN pincer complexes.81 The EPR spectrum of the [Fe(CO)2LPNN](BF4) (LPNN ¼ 2-[(di-tertbutylphosphino)methyl]-6-[1-(2,4,6-mesitylimino)ethyl]pyridine) complex was characterized by a rhombic signal consistent with a S ¼ 1/2 system with g values [2.003, 2.044, 2.088] (giso ¼ 2.045) and observable superhyperfine coupling to the 31P nucleus (I ¼ 1/2). The deviation of the observed g values from the free electron value (ge) as well as the difference between the largest and smallest components of the g tensor (Dg) can be used to estimate the amount of spin density on the metal. The larger deviation of giso from ge and a higher Dg typically correspond to higher spin density on the metal.82 In this study, the large deviation from the ge value as well as the high Dg (¼0.085) suggest that the spin density is mainly located on the iron center. These findings were further supported by MCD measurements and theoretical calculations, resulting in an assignment of an Fe(I) complex. An additional example of the application of EPR spectroscopy in studies of paramagnetic organometallic complexes containing potentially redox non-innocent ligands comes from the work of Dub, Mao, Zhang and coworkers in the studies of organometallic vanadium complexes.83 In this study, the EPR spectra of a (40 -(R)tpy)V(CH2Si(CH3)3)2 (R ¼ CH2Si(CH3)3, C6H5) complexes supported by a terpyridine (tpy) ligand exhibited a rhombic signal characteristic of a S ¼ 3/2 system with an observable metal hyperfine interaction with 51V (I ¼ 7/2). While this signal could originate from a V(III) complex (d2, S ¼ 1) ferromagnetically coupled to the radical ligand or from the V(II) complex (d3, S ¼ 3/2) supported by the neutral ligand, the rhombicity of the signal was attributed to the extended delocalization of unpaired electrons throughout the ligand, suggesting the presence of a ligand radical. Additionally, magnitude of hyperfine coupling as well as observed rhombicity in the hyperfine coupling constant tensor further supported the presence of the ligand based radical, leading to the assignment of the spectra to a V(III) complex ferromagnetically coupled to a radical ligand. It should be noted that perpendicular mode X-band EPR spectroscopy is equally applicable to the study of paramagnetic lanthanide and actinide organometallic complexes.84–89 For example, Arnold, Hohloch and coworkers used X-band EPR spectroscopy for characterization of the (CpiPr4)2U(m-N)-B(C6F5)3 (CpiPr4 ¼ tetra(isopropyl)cyclopentadienyl) complex.90 This complex was described as U(V) with one unpaired electron in the f orbitals, yielding a rhombic EPR spectrum with effective g values of 3.38, 1.01 and 0.79. The large deviation from the ge value is a consequence of the large angular momentum of an unpaired electron in an f orbital. In a further example by Evans and Moehring, EPR spectroscopy has been used for the determination of electron transfer reactivity of rare earth metal complexes.91 Initial studies focused on the reaction of LaIIICp0 3(THF) with [YIICp0 3]− (Cp0 ¼ trimethylsilylcyclopentadienyl). Y(II) is a d1 ion with a S ¼ 1/2 ground state and nuclear spin I ¼ 1/2, while the Ln(III) complex does not have any unpaired electrons and, thus, it is not expected to be observable by EPR. The EPR spectrum of the reaction mixture (Fig. 8) shows the presence of an intense doublet coming from Y(II) due to hyperfine splitting of the I ¼ 1/2 89Y nucleus. In addition, a weaker octet signal can be observed, originating from [LaIICp0 3]−1 formed after reduction of the starting La(III) complex. The La(II) complex features one unpaired electron and, due to the I ¼ 7/2 nuclear spin of La (139La), it is observed in the spectrum as an octet. Thus, this EPR study unambiguously demonstrated that the Y(II) complex can reduce the La(III) complex. Additional reactions with different rare earth metals and ligands were also studied using the same approach, enabling the determination of the reducing abilities of complexes featuring different metals as well as the influence of different ligands on the reducing strengths. Notably, the hyperfine coupling constant, A, increased as the metal character in the singly occupied molecular orbital increased,92 leading the authors to hypothesize that the larger hyperfine coupling might correlate with the increase in the reducing ability of the M(II) complexes. A series of Y and La complexes with known A values was used to test this hypothesis, and a loose correlation between the metal hyperfine coupling constants and the redox potential was observed.

Characterization Methods for Paramagnetic Organometallic Complexes

145

Fig. 8 X-band EPR spectrum of reaction of LaIIICp0 3(THF) complex with the [YIICp0 3]−1 (room-temperature, red) and corresponding simulation (black). Reprinted with permission from Moehring, S. A.; Evans, W. J. Organometallics 2020, 39, 1187–1194. Copyright 2020 American Chemical Society.

For integer spin systems, parallel mode X-band EPR can be used in complexes where the rhombic splitting of the organometallic complex is smaller than the microwave energy (0.3 cm−1).46 This is characteristic of highly axial complexes with small rhombic distortion. In these cases, signals in the parallel mode spectrum are present at lower magnetic fields (e.g., unusually high g values when compared to perpendicular mode). For example, Münck, Holland and coworkers used parallel mode X-band EPR for the characterization of LFeIICH3 (L ¼ a-diketiminate), which is not active for perpendicular mode X-band EPR due to the S ¼ 2 Fe(II) ground state.93 The energy level diagram for a system with S ¼ 2 as a function of rhombicity and the observed parallel mode EPR spectrum are shown on Fig. 9. The broad transition at 59.4 mT is assigned to the quasi-degenerate (D < 0.3 cm−1) Ms ¼  2 doublet. Since the selection rule in parallel mode EPR is DMs ¼ 0, the observed transition should be formally forbidden. However, the rhombic ZFS mixes the Ms ¼ +2 and −2 states, and the transition becomes allowed because it has a DMs ¼ 0 component. The distribution of D values leads to broadness in the observed transition. Spectral simulations suggested that D  0.03 cm−1 and, taking uncertainties in D into account, geff was estimated to be in the range 11.30–11.46. Further analysis of D values lead to the conclusion that investigated complex is most likely in the dz2 ground state (dz2 orbital is lowest in energy). Analysis of geff values along with theoretical calculations were used to determine the energy splitting between d orbitals. Even though it can provide valuable information about integer spin systems, parallel mode X-band EPR is used less nowadays. The reason for this is its limitation to specific integer spin systems and increased availability of high field EPR which can be used for analysis of integer spin systems.

1.06.2.3.2

High field EPR

Limitations of X-band EPR with respect to g value resolution can be overcome by utilizing high frequency (and high field) EPR. With a higher magnetic field, the field dependent interactions are separated from the field independent interactions, which leads to better g value resolution.94 Other advantages of high field EPR (HFEPR) include the ability to investigate integer spin systems, with the appropriate frequency, as well as the determination of ZFS parameters through simulation of experimental spectra.46,95 High field EPR has been successfully utilized for studying organometallic complexes.96–98 The following example demonstrates the advantage of high field over X-band EPR for the determination of ZFS parameters. In the work of Telser, Enders and coworkers, high field EPR was used for the characterization of a series of three Cr(III) complexes with cyclopentadienyl (Cp) ligands (Fig. 10).99 Since these complexes contain Cr(III), a d3 metal, they are also accessible via traditional perpendicular mode X-band EPR. The X-band EPR spectrum for complex 3 exhibits an axial EPR signal typical for a S ¼ 3/2 system where ZFS is higher than the energy of the X-band microwaves. This spectrum also suggested that the axial ZFS parameter, D, is positive such that MS ¼  1/2 doublet of the S ¼ 3/2 ground state is lower in energy than the MS ¼  3/2 doublet. This results in an observed perpendicular feature at g?  4 and a parallel feature at g‖  2. HFEPR was used to obtain the ZFS parameters for all three complexes. The HFEPR spectrum of solid 2 is shown in Fig. 10, along with the simulations with positive and negative D. Comparison of the experimental spectrum to these simulations suggested that the axial ZFS parameter D is positive, with the best fit giving D ¼ + 3.126 cm−1 and E ¼ + 0.087 cm−1 (E/ D ¼ 0.028). In a further example, Marks, Delferro and coworkers demonstrated the use of HFEPR for evaluating the ground state ZFS parameters of the Mn(III) (d4) complex (TMEDA)MnMe3 (TMEDA ¼ N,N,N0 ,N0 -tetramethylethylenediamine).100 While no transitions were observed in perpendicular mode X-band EPR, the HFEPR spectrum of this complex exhibited multiple transitions across a wide field range. Based on the comparison of spectral simulations for both a positive and negative axial ZFS parameter D with the experimental spectrum, the authors determined that D was negative for this complex. The observed pattern in the HFEPR spectrum corresponded well to a quintet ground state with D ¼ − 2.178 cm−1. Based on these results, the authors concluded that the electronic structure of this complex arises primarily from the elongation of the axial Mn–N bond. This distinguishes the x and y axes from the z axis and leads to the negative D value which corresponds to a dx2−y2 ground state (dx2−y2 is unoccupied and the highest energy d orbital, while rest of the d orbitals are partially occupied). Lastly, the value and sign of D are similar to the Mn(III) tetrapyrrole complexes which adopt square planar/pyramidal symmetry,95,101–103 suggesting that (TMEDA)MnMe3 also has a square-pyramidal geometry.

146

Characterization Methods for Paramagnetic Organometallic Complexes

Fig. 9 (A) Splitting of the energy levels for S ¼ 2 system as a function of rhombicity (D ¼ 3D(E/D)2). (B) Parallel mode X-band EPR spectrum of LFeIICH3 (2 K). Adapted with permission from Andres, H.; Bominaar, E. L.; Smith, J. M.; Eckert, N. A.; Holland, P. L.; Münck, E. J. Am. Chem. Soc. 2002, 124, 3012–3025. Copyright 2002 American Chemical Society.

Fig. 10 (A) Structures of studied Cr complexes. (B) X-band EPR spectrum of complex 3 in dichloromethane (4.8 K, black) with corresponding simulation (red). (C) HFEPR spectrum of complex 2 (7 K and 321.6 GHz, black) and corresponding simulations with positive D (red) and negative D (blue). Adapted with permission from Krzystek, J.; Kohl, G.; Hansen, H.-B.; Enders, M.; Telser, J. Organometallics 2019, 38, 2179–2188. Copyright 2019 American Chemical Society.

Characterization Methods for Paramagnetic Organometallic Complexes

1.06.2.4

147

Pulsed EPR spectroscopy

To evaluate weak interactions, such as hyperfine interactions with ligand atoms that are not directly bound to the metal, pulsed EPR techniques must be used.36 In pulsed EPR, short microwave pulses (10–100 ns) are used to excite the sample and transient emission from the sample is measured following excitation. Since the pulses are short, they have a limited excitation energy range and often cannot excite the whole spectrum.104 The frequencies used in pulsed EPR are the same as those employed in CW EPR but a different detection system is used. While the experimental system is more complicated, it provides the possibility to change more parameters, such as magnetic field, number of pulses, time delays, pulse lengths and frequencies. Pulsed EPR is especially effective for the determination of weak interactions such as dipole-dipole interactions between electron spins or hyperfine couplings, as well as for the resolution of complicated spectra and the measurement of relaxation times. The most basic pulsed EPR experiment uses one pulse and measures the amplitude of free induction decay (FID) as a function of magnetic field where the intensity is highest if the pulse is p/2. In the two-pulse experiment, the primary echo is measured as a function of magnetic field. The pulse sequence is p/2–t–p and the primary echo is measured after time t. This pulse sequence is the basis for two pulse electron spin echo envelope modulation (ESEEM), echo detected Davis electron-nuclear double resonance (ENDOR) and relaxation time measurements. In the three-pulse experiment, the pulse sequence is p/2–t–p/2–T–p/2 and the stimulated echo is measured after time t. The three-pulse sequence is the basis for Mims ENDOR, three pulse ESEEM and hyperfine sub-level correlation (HYSCORE) EPR.104–108 The samples for pulsed EPR methods are most commonly in the solid state (crystals, powders, frozen solution) because only in solid samples are the relaxation times of metal systems long enough to allow application of the pulses. Also, lowering the measurement temperature can decrease relaxation times. Commonly, the temperature utilized in pulsed EPR is 10 K or lower for high-spin systems and metal ion clusters or 10–30 K for S ¼ 1/2 systems. For frozen solution samples, solvents need to be chosen to give clear optical glasses upon freezing in order to prevent aggregation of the sample. Frequency, repetition rate, and pulse lengths all depend on the properties of the sample and chosen method.104 Pulsed EPR methods have also been successfully used for studying organometallic systems.109–112 An example of the utility of pulsed EPR techniques in organometallic chemistry is from Mills and coworkers to evaluate covalency in actinide complexes, An(Cptt)3 (An ¼ Th (1), U (2); Cptt ¼ 1,4-di-tert-butylcyclopentadienyl) (Fig. 11).113 Covalency in such systems can be measured via the superhyperfine interaction of the metal based unpaired electrons with nuclear spins in surrounding ligands. Since this interaction is often not resolved in CW EPR, pulsed methods are essential. In this example, the authors used ESEEM (1D technique) and HYSCORE (2D ESEEM technique) EPR to measure electron spin densities at the 13C and 1H nuclei of ligands in both Th and U complexes. The CW X-band EPR of the thorium compound confirmed that the Th ion had an electron configuration of 6d15f0, which gives an axial EPR spectrum with g‖ ¼ 1.974 and g? ¼ 1.880 (consistent with an unpaired electron in dz2 orbital). The X-band EPR spectrum of compound 2 was consistent with a 5f3 ground state electronic configuration resulting in an EPR spectrum with gx ¼ 3.05, gy ¼ 1.65 and gz < 0.5 (where the latter is not observable within the magnetic field range utilized in this study). These electronic structures were in accordance with additional magnetic studies and theoretical calculations. Both complexes exhibited high modulation amplitudes in ESEEM measurements due to the interaction with the 1H on the ligands. In order to quantify these interactions, HYSCORE was used. Fig. 12 shows the HYSCORE spectra of Th complex in both the 13C and 1H regions. In HYSCORE, the magnetization components that are undisturbed during the pulse sequence give rise to diagonal peaks in the 2D spectrum. The magnetization components that are exchanged and/or transferred during the pulse sequence give cross-peaks, which are symmetric across the diagonal. For weak coupling where the nuclear Zeeman interaction is stronger than hyperfine interaction (2| nI | > A), it is expected that peaks are symmetrically centered around the Larmor frequency of the nuclei (nI) and separated by the hyperfine coupling constant. For complex 1, the 13C region reports the p spin density in the frontier orbitals of ligands. Two distinct peaks can be observed in this region, one on the antidiagonal and another wider peak next to it. This suggests that there are at least two different 13C positions in this complex. Through the data analysis, values of hyperfine couplings were obtained (for C2 A‖,? ¼ +3.7, +0.4 MHz, and for C1,3 A‖,? ¼ + 1.1, +0.4 MHz) and it was calculated that the 2pp spin populations of C2 and C1,3 are 1.3% and 0.5%, respectively. This was supported by an analysis of the 1H region, which showed that the 2pp spin populations at C4,5 are negligible. Summing the 2pp spin populations across all three ligands leads to a total spin population of 6% of spin on the

Fig. 11 Structures of investigated actinide complexes and numbering scheme for ligands.

148

Characterization Methods for Paramagnetic Organometallic Complexes

Fig. 12 HYSCORE spectra of Th complex in toluene (11 K) at (A) 13C region (B0 ¼ 366.3 mT) with simulation (red), (B) 1H region (B0 ¼ 351.6 mT). Adapted with permission from Formanuik, A.; Ariciu, A. M.; Ortu, F.; Beekmeyer, R.; Kerridge, A.; Tuna, F.; McInnes, E. J. L.; Mills, D. P. Nat. Chem. 2017, 9, 578–583. Copyright Springer Nature 2017.

ligands. A similar analysis was performed for complex 2, yielding a total spin population of 17% on the ligands. The larger spin density on the ligands for the uranium complex compared to the thorium analog was surprising considering higher radial extent of 6d orbitals in 1 versus 7f orbitals in 2. However, this suggests that the angular parts of the orbitals have a significant influence on bonding. In the case of thorium complex, only the 6dz2 orbital has the correct orientation to overlap with frontier ligand orbitals, while the three singly occupied 5f orbitals in uranium complex have greater overlap with the ligand orbitals, resulting in higher spin population.

1.06.3

Magnetic circular dichroism (MCD) spectroscopy

1.06.3.1

Introduction

MCD spectroscopy combines the circular dichroism (CD) technique with an applied longitudinal magnetic field, measuring the differential absorption of left and right circularly polarized light. While in traditional CD spectroscopy only optically active molecules can be studied, in MCD the applied magnetic field induces optical activity via the Faraday effect, enabling studies on molecules without inherent optical activity. While MCD has historically been more extensively employed in bioinorganic chemistry, particularly studies of active sites of metalloproteins, it has recently become more widely utilized for studying paramagnetic organometallic compounds due to its advantages over other spectroscopic techniques. For example, the dependence on the magnitude of SOC on signal intensity in MCD results in a significant increase in intensity for metal-based transitions due to the large SOC constants of metals. This results in an increase in the intensity for metal-based transitions such as d-d and f-f transitions, which are Laporte forbidden and often difficult to observe in electronic absorption spectroscopy. Direct observation of these transitions can provide information about ligand field strength in the organometallic complexes. MCD can also provide detailed information on the electronic structure and geometry of organometallic complexes by probing excited state transitions as well as insight into the ground state of a paramagnetic organometallic species (via saturation magnetization experiments). Furthermore, in combination with freeze-trapped methods, MCD can be used for the identification and characterization of organometallic intermediates in different chemical transformations. Additionally, MCD can be used to study the magnetization of one species at a time, unlike bulk magnetization techniques (e.g., SQUID), which is useful in analysis of systems that contain more than one species (e.g., catalytic reactions). Importantly, this method can be used for any S > 0 organometallic system, including both Kramers and non-Kramers ground states. Beside above mentioned advantages of MCD, its use remains limited due to the necessity for superconducting magnets, cryogenic temperatures for C-term MCD, and glassing solvents for frozen solution samples.

1.06.3.2

Theory

The MCD experiment is built upon traditional CD spectroscopy. In the CD experiment, electronic transitions must be both electric dipole and magnetic dipole allowed, which is satisfied for chiral molecules which have a helical distribution of charge. Based on this helical electronic charge distribution, chiral molecules preferentially absorb either left circularly polarized (LCP) or right circularly polarized (RCP) light. This difference in absorption of LCP and RCP light is measured as the ellipticity DA. In MCD, an applied longitudinal magnetic field removes the degeneracy of states according to the Zeeman effect, and different ML states have different

Characterization Methods for Paramagnetic Organometallic Complexes

149

LCP and RCP radiation absorption. The selection rule in MCD spectroscopy, DML ¼  1, requires the presence of orbitally degenerate ground or excited states in order to have MCD transitions. As the ground and excited states are not necessarily degenerate, the degeneracy requirement can be satisfied through spin-orbit coupling (SOC), which is significant in organometallic complexes as metals have large SOC constants. The intensity of an MCD spectrum is a function of three terms, called the A-, B- and C-term:         DA 2N0 p3 a2 Cl log e ∂f ðEÞ C0 ¼ mB H A1 − + B0 + f ðEÞ (8) E 250hcn ∂E kT where DA is the field-dependent difference between LCP and RCP light, a is the electric permeability, C is the concentration, l is the path length, n is the index of refraction, mB is the electron Bohr magneton, H is the applied magnetic field, and f(E) is the absorption band shape.21 In this expression, A1 represents the A Faraday term which arises from Zeeman splitting of orbitally degenerate ground and/or excited states. It has a derivative shape due to the separation of the individual LCP and RCP radiation absorbing bands. This term is not temperature dependent, nor is the B-term (B0), which arises due to the field induced mixing of states. The C-term (C0, Fig. 13) requires the presence of a degenerate ground state. This term is dominant at low temperatures (kT  gmBH) and has an absorption band shape. C-term MCD is the most important of the three intensity mechanisms when studying paramagnetic organometallic complexes. For most open shell organometallic complexes, the ground state is generally orbitally nondegenerate, which means that the degeneracy is achieved through SOC. In the case where a complex has degenerate excited states, it is possible to observe temperature dependent pseudo-A terms, which have a derivative shape that is produced by two C-terms of opposite sign, with an energy difference that is proportional to the SOC in the excited state. Importantly, the intensity of the C-term MCD signal also depends on the magnitude of SOC, which means that the intensity of the metal-centered transitions, for which SOC is larger than for ligand-centered transitions, is disproportionately high relative to the intensity of that transition in electronic absorption spectra.21,23,114 In addition to excited state information, C-term MCD can provide significant insight into the ground states of paramagnetic organometallic complexes, including g values and ZFS parameters, via analysis of the temperature and magnetic field dependence of the C-term signal. This method is often referred to as variable temperature variable field (VTVH) MCD or saturation magnetization MCD.115,116 Information obtained by this method are analogous to the information that can be obtained by bulk magnetization techniques like VTVH SQUID, but as mentioned above, a major advantage is that VTVH MCD can provide information about individual species within a mixture. The example of the signal intensity dependence on the applied magnetic field for a S ¼ 1/2 system is shown on Fig. 14. At low magnetic fields, where Zeeman splitting is small, both ground state sublevels are equally populated. This means that the absorbance of the LCP and RCP radiation is equal and the observed field dependence on MCD intensity is linear. As the field increases, the sublevels become more split and the lower ground state sublevel becomes more populated due to the Boltzmann distribution. This results in the preferential absorption of RCP radiation which leads to the saturation of the signal. The saturation curves are measured at a several different temperatures and can be used to determine the spin of the ground state. Detailed fitting and analysis also enable further information about the ground state to be determined including g values, ZFS parameters and, with Kramers ground state systems where the g values are experimentally determined, the polarizations of transitions even in non-crystalline samples.

Fig.13 Energy level diagram showing absorption of LCP and RCP for degenerate ground state (C-term) and corresponding Faraday C-term peak shape.

150

Characterization Methods for Paramagnetic Organometallic Complexes

Fig. 14 Saturation magnetization curve for S ¼ 1/2 system.

1.06.3.3

Applications

As base metals continue to attract more interest in organometallic chemistry and catalysis, methods such as MCD spectroscopy are critical to providing insight into electronic structure as paramagnetism is much more prevalent with these metals compared to precious metal systems. For example, near infrared (NIR) MCD enables the direct observation of d-d transitions, as these transitions are often too low in energy to see in visible spectra, while UV-Vis MCD can provide high resolution spectra (better than absorption spectroscopy) to identify and assign charge transfer (CT) transitions. The combination of experimental MCD spectra with theoretical calculations can provide further insight into electronic structure. In addition, VTVH MCD can provide information about spin and oxidation state, as well as ZFS parameters for the ground state. Beyond that, in combination with freeze-trapping, MCD can be used for identification and characterization of in situ formed organometallic species. The following examples highlight recent applications of MCD for studies of paramagnetic organometallic complexes, demonstrating the utility of this method. MCD was recently used for comparing electronic structure and bonding in four-coordinate iron N-heterocyclic carbene (NHC) complexes to amine and phosphine systems, for insight into the distinct catalytic abilities of iron-NHCs.117–121 Using MCD spectroscopy complemented by DFT and 57Fe Mössbauer investigations, Neidig and coworkers obtained insight into ligand field effects and bonding across these three ligand types.122 To demonstrate how MCD can be utilized in such studies, (IMes)2FeCl2 (IMes ¼ 1,3-bis(2,4,6-trimethylphenyl)imidazol-2-ylidene) serves as a detailed example (Fig. 15A). While in ideal tetrahedral symmetry only one spin allowed ligand field transition (5E ! 5T2) would be predicted for this complex, its NIR MCD spectrum (Fig. 15B) shows two distinct but energetically close transitions. This is consistent with the lower symmetry of this molecule, where the degeneracy of the ground and excited states is removed. NIR MCD was also used to determine the ligand field strength of this complex. In order to determine ground spin state and ZFS parameters, VTVH MCD was used (Fig. 15B). Fitting saturation magnetization data obtained at 5917 cm−1 indicated a high spin Fe(II) (S ¼ 2) with negative ZFS, and all ground state spin Hamiltonian parameters were determined. The UV-Vis MCD spectrum shows that there are multiple CT bands (Fig. 15C). These bands, as well as ligand field transitions, could be assigned using time dependent DFT calculations considering only transitions with significant metal character that show significant C-term MCD intensity. To evaluate relative ligand field strengths of the NHC ligands compared to diamine and phosphine ligands, the NIR MCD studies discussed above for (IMes)2FeCl2 were expanded to additional four-coordinate, bis-chloride NHC, amine and phosphine complexes. Of note, the spectra of the diamine ligand containing complexes were like those of complexes with NHC ligands with two positive bands. In contrast, the spectra of complexes with phosphine ligands also shows two bands, but one is positive and the other is negative. This is a consequence of a pseudo-A term in the ligand field MCD transitions of the phosphine complex which arises from spin-orbit coupling between two energetically close excited states. Overall, these experiments determined that the ligand field splitting of the NHCs was larger than those of the amine complexes but smaller than those of the phosphine systems. In iron(II)-NHC complexes featuring saturated, unsaturated and unsaturated/chlorinated NHC ligands, the d-d transitions in the NIR MCD and the saturation magnetization data were similar, indicating that the backbone substitution had little effect on the ligand field strength of the iron-NHC complexes. The combined effect of backbone and N-substituents was studied in more detail by MCD for alkylated iron(II) NHC complexes.123 The investigated complexes included (NHC)Fe(1,3-dioxan-2-ylethyl)2 where NHC ¼ IMes, IPr (1,3-bis(2,6-diisopropylphenyl)imidazol-2-ylidene) and their saturated versions SIMes and SIPr. The NIR MCD spectra suggested that all of these complexes were high spin iron(II). The NIR spectra of the complexes bearing IMes and SIMes ligands revealed two negative features (at 5310

Characterization Methods for Paramagnetic Organometallic Complexes

151

Fig.15 (A) Structure of (IMes)2FeCl2 complex. (B) NIR and (C) UV-Vis MCD spectrum of (IMes)2FeCl2 (5 K and 7 T). Insert shows VTVH MCD data collected at 5917 cm−1. Adapted from Fillman, K. L.; Przyojski, J. A.; Al-Afyouni, M. H.; Tonzetich, Z. J.; Neidig, M. L. Chem. Sci. 2015, 6, 1178–1188, with permission from the Royal Society of Chemistry.

and 7600 cm−1 for IMes complex, and 5400 and 7630 cm−1 for SIMes complex) and showed that there is small difference in ligand field strength for these two complexes (6455 and 6515 cm−1 for IMes and SIMes, respectively). This demonstrated that there is no significant difference in the electronic structure as a result of NHC saturation, which is in accordance with the previously described study.122 However, for the complex with the SIPr ligand, the combined steric effects of both the saturated backbone and N-substituent prevented chelation of two alkyl oxygens, resulting in a significantly perturbed electronic structure as observed in the NIR MCD spectra. Spectra revealed presence of two C-term transitions for IPr complex, and one pseudo A-term for SIPr complex. It was also determined that ligand field strength for SIPr complex (7140 cm−1) is significantly higher than for IPr complex (6530 cm−1). The shift in ligand field strength relates to change in iron site geometry going from pseudo five-coordinate (in IPr complex) to pseudo four-coordinate (SIPr complex). Subsequent studies by Neidig and coworkers extended the use of MCD for studies of iron(II) pincer NHC complexes, focusing on the influence of NHC moieties in pincer ligands on electronic structure and bonding in (pincer)FeBr2 complexes (Fig. 16).124

152

Characterization Methods for Paramagnetic Organometallic Complexes

Fig. 16 Near IR MCD spectra (5 K and 7 T) and structures of (A) (iPrCDA)FeBr2, (B) (iPrPDI)FeBr2 and (C) (iPrCNC)FeBr2 complexes. Adapted with permission from Baker, T. M.; Mako, T. L.; Vasilopoulos, A.; Li, B.; Byers, J. A.; Neidig, M. L. Organometallics 2016, 35, 3692–3700. Copyright 2016 American Chemical Society.

NIR and VTVH MCD studies were used to determine spin state, ligand field strengths and ZFS parameters in these systems. The NIR MCD spectrum of (iPrCDA)FeBr2 (Fig. 16A) contained five ligand field (LF) transitions, which is inconsistent with a high spin Fe(II) and consistent with intermediate spin (S ¼ 1) Fe(II). The S ¼ 1 spin state was confirmed by VTVH MCD measurements. In contrast, the (iPrPDI)FeBr2 and (iPrCNC)FeBr2 (Fig. 16B and C) complexes exhibited two LF transitions in NIR MCD spectra, consistent with a high spin Fe(II) which was also confirmed by VTVH MCD. In (iPrCNC)FeBr2, the introduction of two NHC ligands leads to an increased ligand field due to the increase in s donation when the NHC ligands replace imine nitrogens in the corresponding (iPrPDI)FeBr2 complex. However, while the ligand field strength is higher in (iPrCNC)FeBr2, it is still S ¼ 2 in contrast to (iPrCDA) FeBr2 where the ligand field strength is high enough to give an S ¼ 1 ground state. This indicates that the position of NHC ligands is a significant factor in the ligand field strengths and ground state spins as opposed to simply the number of NHC ligands present in a system. C-term MCD spectroscopy has also been employed to evaluate the electronic structure and bonding in cobalt NHC complexes.125 The MCD spectrum of (IMes)2CoCl2 features six LF transitions, three low-energy and three high-energy transitions (Fig. 17). This is consistent with the distorted tetrahedral high spin Co(II) (S ¼ 3/2), which has a characteristic spectrum with strong negative signal and series of positive signals at higher energies, based on previous studies.126,127 The two groups of transitions are derived from the spin allowed 4A2 ! 4T1(F) and 4A2 ! 4T1(P) transitions in tetrahedral symmetry which, due to the lower symmetry of this complex, further split to give six total transitions. MCD was also used in this study to determine the ligand field strength in (ICy)2CoCl2 (ICy ¼ 1,3-dicyclohexylimidazol-2-ylidene) and (dppp)CoCl2 (dppp ¼ 1,3-bis(diphenylphosphino)propane) complexes in order to compare ligand field strengths. The MCD spectra of all these complexes contain the same general features, since they are all distorted tetrahedral high spin Co(II) complexes. Comparing the NIR MCD spectra of (ICy)2CoCl2 and (IMes)2CoCl2 in more detail revealed that the lowest energy transitions are shifted toward higher energy (1000 cm−1) while the higher energy transitions are slightly shifted toward lower energy. This suggested that the N-substituent of the NHC can have a significant effect on the metal NHC bonding, likely due to the steric effects of the substituent. By contrast, for (dppp)CoCl2 each of the NIR transitions are shifted toward higher energy, consistent with a stronger ligand field than present in the Co(II)-NHC complexes.

Characterization Methods for Paramagnetic Organometallic Complexes

153

Fig. 17 MCD spectrum of (IMes)2CoCl2 (5 K and 7 T). Adapted from Iannuzzi, T. E.; Gao, Y.; Baker, T. M.; Deng, L.; Neidig, M. L. Dalton Trans. 2017, 46, 13290–13299, with permission from the Royal Society of Chemistry.

MCD has also proven to be a valuable technique to evaluate electronic structure and bonding in paramagnetic f-element organometallic complexes. C-term MCD spectroscopy is a suitable technique for studying these complexes because it enables the observation and increased resolution of f-f transitions, which can be more challenging to resolve by traditional room temperature UV-Vis absorption spectroscopy. As a representative example, C-term MCD spectroscopy was recently utilized to evaluate electronic structure in a series of lanthanide(II) complexes, [Ln(Cp0 )3]− (Cp0 ¼ trimethylsilylcyclopentadienyl) (Fig. 18A), which were synthesized via reduction of the corresponding lanthanide(III) complexes.128 Previous spectroscopic data suggested that reduction would add an electron to the d orbital rather than the f orbital. MCD spectroscopy was used to confirm these findings as well as to provide further insight into the electronic structure of these complexes. The complexes of La (5d1), Pr (4f2 5d1), Gd (4f7 5d1), and Eu (4f7), as well as Y (4d1) as a d-block analog were characterized by both NIR and UV-Vis MCD. Only the NIR spectrum of the Pr complex contained MCD-observable f-f transitions (Fig. 18B), as only the Pr complex has two electrons in f orbitals while the other complexes have unoccupied or half-occupied f orbitals. The observed bands for the Pr complex can be fit to at least five different transitions. These transitions are also broader (fwhm 600 cm−1) than expected for pure f-f transitions (fwhm 50 cm−1), but less broad than expected for d-d transitions (fwhm 1000 cm−1). This suggested that the ground state molecular orbitals are mixtures of f and d orbitals. Lastly, each of the complexes showed multiple intense transitions in the UV-Vis region (Fig. 18C) and, due to the higher resolution of MCD spectroscopy, more transitions could be resolved than in previous UV-Vis absorption experiments. Using theoretical calculations, the lowest energy transitions could be assigned to electric dipole-allowed ndz2 ! diffuse (n + 1) pz transitions, followed by dz2 ! quasi-degenerate (n + 1)px/py transitions. An additional example of the use of MCD spectroscopy for the characterization of paramagnetic organometallic actinide complexes comes from a recent study of homoleptic uranium (IV) aryl complexes by Neidig and coworkers.129 The six-coordinate U(IV) complexes, [U(Ar)6]2− (Ar ¼ Ph (1), p-tolyl (2), p-Cl-Ph (3)), were synthesized and MCD spectroscopy was used to provide further insight into the effect of aryl ligand donor ability on electronic structure in these complexes. The NIR MCD spectra of these complexes contained many f-f transitions (Fig. 19). The NIR spectra for 1 and 2 are almost identical, and all transitions can be grouped to three energy regions 5000–7000 cm−1 (features I), 8000–9500 cm−1 (features II), and 10,700 cm−1 (features III). The spectrum for complex 3 is slightly different in terms of energy and number of transitions, but these differences are minimal and NIR spectra of the three complexes supports similar overall electronic structures and, hence, minimal aryl effects. Beyond studies of electronic structure and bonding in well-defined, isolated complexes, MCD spectroscopy in combination with freeze-quench techniques can be a powerful method for identification and characterization of in situ formed organometallic species as well as intermediates in catalysis. One example of applying MCD to mechanistic studies focused on the iron-catalyzed Kumada cross-coupling of MesMgBr and primary alkyl halides using the bisphosphine ligand, SciOPP (1,2-bis[bis{3,5-di(t-butyl)phenyl} phosphino]benzene).130 This reaction is catalyzed by mesityl-Fe(II)-SciOPP complexes formed in situ from (SciOPP)FeCl2 and MesMgBr. Identification of the in situ formed iron species, with both stoichiometric and excess amounts of MesMgBr, was accomplished using MCD and 57Fe Mössbauer spectroscopies. The MCD spectrum from addition of 1 equivalent of MesMgBr contained two LF transitions (Fig. 20A), characteristic of a high spin Fe(II) distorted tetrahedral complex. This was assigned to (SciOPP)FeBrMes complex via comparison to the spectrum for isolated crystalline material. Addition of two equivalents of Grignard reagent resulted in a significant change in the NIR MCD spectrum (Fig. 20B). The spectrum contained multiple low-energy LF transitions and could be assigned to the distorted square planar S ¼ 1 (SciOPP)FeMes2 complex which was also isolable. Further increasing the amount of MesMgBr to 20 and 100 equivalents changed the NIR MCD spectrum yet again (Fig. 20C and D). These new features were assigned to FeMes−3 by comparison of the MCD spectra of the independently synthesized compound with the MCD spectrum of the in situ formed species. Further MCD was also used for characterization of the reaction of (SciOPP)FeMes2 and FeMes−3 with primary alkyl halide, demonstrating the consumption of (SciOPP)FeMes2 and formation of a high-spin (SciOPP) FeXMes product (X ¼ halide). A similar approach was also used in studies of iron− SciOPP catalyzed Suzuki −Miyaura and Kumada cross-coupling of phenyl nucleophiles and secondary alkyl halides.131 In these studies, MCD spectroscopy enabled identification of

154

Characterization Methods for Paramagnetic Organometallic Complexes

Fig. 18 (A) Structure of studied [Ln(Cp0 )3]− complexes. (B) NIR MCD spectrum of [Pr(Cp0 )3]− (5 K and 7 T). (C) UV-Vis spectra of [Y(Cp0 )3]−, [Pr(Cp0 )3]− and [Gd(Cp0 )3]− (5 K and 7 T). Adapted with permission from Fleischauer, V. E.; Ganguly, G.; Woen, D. H.; Wolford, N. J.; Evans, W. J.; Autschbach, J.; Neidig, M. L. Organometallics 2019, 38, 3124–3131. Copyright 2019 American Chemical Society.

the two in situ formed bis-phenylated Fe(II)-SciOPP species observed by Mössbauer spectroscopy. The NIR MCD spectrum of a frozen solution of these complexes contained four LF transitions at 5600, 6560, 8480, and 13,500 cm−1, indicative of the presence of at least two Fe(II) species. The low energy transitions are consistent with a distorted tetrahedral high spin Fe(II) compound, while the band at 13,500 cm−1 was too high in energy for a four-coordinate, tetrahedral species, and instead was indicative of a five-coordinated distorted square-pyramidal high spin Fe(II) compound. The assignment of two different high-spin iron(II) species was further confirmed by VTVH MCD. The four-coordinate tetrahedral species was assigned to diphenyl complex ((SciOPP)Fe(Ph)2), while five-coordinate distorted square pyramidal species was assigned to the THF adduct ((SciOPP)Fe(Ph)2(THF)). Beyond iron catalysis, MCD spectroscopy has also been used to characterize high-valent Ni species in C-C cross coupling reactions. For such systems, the reductive elimination (RE) step is rate determining and the observed RE rates in high-valent Ni complexes vary by over five orders of magnitude, depending on the metal’s oxidation state and ligands.132,133 Park and coworkers used MCD along with other spectroscopic techniques to determine RE rates and elucidate electronic structures of cycloneophyl Ni(III) and Ni(IV) complexes (Fig. 21C).134 Since a crystal structure was only available for complex 4a, the MCD and EPR were used to elucidate the structures of the other high-valent Ni complexes in this study. The six coordinate Ni(III) complex 4a features two NIR MCD transitions at 12,300 and 15,300 cm−1, while for complex 3a these transitions are shifted toward lower energy (Fig. 21A). This shift in energy, combined with data from EPR studies, suggests that the 3dz2 orbital is lower in energy for 3a than in 4a. The higher energy of 3dz2 in 4a is attributed to strong s donation of the CF3 ligand. Under noncryogenic conditions where RE occurs, the lowest energy d-d transition in 3a is shifted to 8300 cm−1. This was also observed when using non-coordinating solvents, suggesting that under the RE reaction conditions solvent is not coordinated. Thus, it was concluded that 3a is a five-coordinate complex with square pyramidal geometry where the pyridine of the py3CH ligand is on the z axis. For the 3b and 4b complexes, where the py3CH ligand is replaced by a bpy (2,20 -bipyridine) ligand, the NIR MCD spectra were more complicated and assigned to a mixture of fiveand six-coordinate complexes, though under RE reaction conditions both complexes have five-coordinate square pyramidal geometries (Fig. 21B). Combining the spectroscopic insight with the previously determined RE activity, it is determined that five-coordinate high-spin complexes are more active for reductive elimination than the six-coordinate complexes. The authors showed that the absorption maxima of the investigated complexes correlate with the RE activity. To gain more insight into the nature of these transitions, the authors utilized the higher resolution of MCD along with theoretical calculations. The results revealed that in all complexes the maxima in absorption spectra corresponds to the charge transfer between 2p s bonding orbital of the coordinated C (from cycloneophyl ligand) to the Ni 3d s antibonding orbital. Based on these findings the authors suggested that the energy of the C-to-Ni charge transfer (LMCT) transition correlates with the RE rate and can be used to predict relative RE activity.

Characterization Methods for Paramagnetic Organometallic Complexes

155

Fig. 19 Near IR MCD spectra of U complexes 1, 2 and 3 (5 K and 7 T). Adapted with permission from Wolford, N. J.; Sergentu, D.; Brennessel, W. W.; Autschbach, J.; Neidig, M. L. Angew. Chem. Int. Ed. 2019, 58, 10266–10270. Copyright 2019 Wiley-VCH Verlag GmbH & Co. KGaA.

1.06.4

X-ray absorption spectroscopy

1.06.4.1

Introduction

X-ray absorption spectroscopy (XAS) is based on the absorption of X-ray radiation by a specific element in the sample, which promotes an electron from the core orbitals to the lowest unoccupied orbitals or to the continuum (ionization). In contrast to EPR and C-term MCD spectroscopies, this technique is not limited to paramagnetic samples and is equally useful for the characterization of organometallic complexes in any spin state. However, it has become one of the most useful methods for studies of paramagnetic organometallic complexes including intermediates, because these more often have ambiguous oxidation states and electronic structures. In paramagnetic organometallic chemistry, XAS is widely used for determining the oxidation state of the metal centers as well as for insight into geometry, electronic structure and covalency. Furthermore, it can be used to study local structure and bonding through the determination of the number, type, and distance of neighboring atoms to a metal center of interest. Since XAS is an element-specific method, open shell organometallic complexes can be studied from the perspective of both the metal centers and the ligands, providing detailed insight into electronic structure and bonding in these complexes. The need for tunable energy X-ray radiation requires the use of synchrotron radiation as an excitation source, which is a limiting factor in terms of the access and use of this method. Even with this challenge, the use of XAS in the characterization of paramagnetic organometallic complexes continues to expand, placing it as one of the go-to methods for advanced characterization of such species.

156

Characterization Methods for Paramagnetic Organometallic Complexes

Fig. 20 NIR MCD spectra from reaction of FeCl2(SciOPP) with (A) 1 equivalent, (B) 2 equivalents, (C) 20 equivalents and (D) 100 equivalents of MesMgBr (5 K and 7 T). Adapted with permission from Daifuku, S. L.; Al-Afyouni, M. H.; Snyder, B. E. R.; Kneebone, J. L.; Neidig, M. L. J. Am. Chem. Soc. 2014, 136, 9132–9143. Copyright 2014 American Chemical Society.

1.06.4.2

Theory

The energy of X-rays is sufficient to eject core electrons from an atom. The core shells have specific binding energies and the X-ray absorption spectrum of any element is characterized by sharp increases in absorption at the X-ray energies corresponding to these binding energies.135 The sharp increases in absorption are called absorption edges and are named according to the principal quantum number of the electron that is excited (K-edge for n ¼ 1, L-edge for n ¼ 2, M-edge for n ¼ 3, . . .). The edge energies for core orbitals increase as the atomic number of an element increases and the Coulomb attraction to the nucleus increases. For K-edge, these edge energies are in the range from 284 eV for carbon to 115,606 eV for uranium, while the corresponding L-edge energies for these elements are much lower.135,136 The energy range in spectrum near the absorption edge is referred to as the X-ray absorption near edge structure (XANES) region of the XAS spectrum. At higher energies an electron is ejected to the continuum and spectral region corresponding to these energies, which can extend for 1000 eV or more, is defined as the extended X-ray absorption fine structure (EXAFS) region (Fig. 22).31,135–139 The XANES region can provide valuable information about oxidation state and coordination environment of an element of interest.31,136,137,140 For organometallic complexes, the metal K-edge is a particularly useful transition, corresponding to the excitation of a core 1s electron to valence bound states localized on the metal or to the continuum (Fig. 22). The edge energy, which is determined via the first inflection point at the absorption edge (or the maximum of its first derivative), can be used to assign the formal oxidation state of a metal center. Generally, oxidation of the absorbing metal increases its electron affinity (binding energy), resulting in a shift of the absorption edge toward higher energy.31,136 The edge energy correlates well with the formal oxidation state in situations with hard ligands and little covalency. In systems with higher covalency this correlation is not as

Characterization Methods for Paramagnetic Organometallic Complexes

157

Fig. 21 (A) Near IR MCD (5 K and 7 T) spectra of (i) cryoreduced 4a, (ii) 3a frozen in PrCN, and (iii) 3a frozen in toluene (blue dashed curve is the 298 K, 5 T MCD spectrum of 3a dissolved in the nitrile solvent). (B) Near IR MCD (5 K and 7 T) spectra of (i) cryoreduced 4b, (ii) 3b frozen in PrCN, and (iii) 3b frozen in toluene (blue dashed curve is the 243 K, 5 T MCD spectrum of 3b dissolved in PrCN). (C) Structures of the studied Ni complexes. Adapted with permission from Shin, J.; Gwon, S.; Kim, S.; Lee, J.; Park, K. J. Am. Chem. Soc. 2020, 142, 4173–4183. Copyright 2020 American Chemical Society.

Fig. 22 Energy level diagram showing K-edge XAS transitions for transition metal complexes with corresponding spectrum.

consistent, and additional attention is necessary for assigning oxidation states. An additional problem arises when the fine structure is overlaid on the absorption edge which can influence one’s ability to define a specific energy of the edge. For first row transition metal complexes the edge corresponds to the 1s to 4p and to 4p + ligand to metal charge transfer (LMCT) transitions (shakedown transitions), and it represents the threshold for ionization.137,141 Furthermore, additional lower energy features can be observed at energies just before the edge in transition metal complexes with unoccupied or partially occupied d orbitals and can provide information about coordination environment. These pre-edge features are superimposed on the rising edge and originate from the 1s to 3d transitions which are electric dipole forbidden, but quadrupole allowed. Since these are dipole forbidden transitions, their intensity is much lower than the rising edge, but the pre-edge transitions can gain intensity through the mixing of 3d and 4p orbitals of suitable symmetry. For centrosymmetric complexes these transitions are weak while for non-centrosymmetric complexes they have significantly higher intensities, due to increased mixing of 3d and 4p orbitals.31

158

Characterization Methods for Paramagnetic Organometallic Complexes

While metal K-edge XAS is most commonly used for the characterization of metal sites in organometallic complexes containing 3d and 4d metals, for heavier transition metals as well as for f-elements, the lower energy L-edge XAS transition is commonly studied, as the K-edge energy is too high.138 In the L-edge transition, core electrons from the second shell are excited including electric dipole allowed transitions from 2p to (n-1)d orbitals, which are more intense than pre-edge transitions in K-edge XAS and thus suitable for probing (n-1)d manifolds. L-edge XAS displays three distinct L-edges. L1 originates from excitation of a 2s electron but, since transition to an (n-1)d orbital is electron dipole forbidden and thus has a weak intensity, this edge is not commonly used. The remaining two L-edges correspond to the excitation of a 2p electron. After excitation, due to the spin orbit coupling in the 2p5 configuration, the excited state is split into 2p3/2 and 2p1/2 giving rise to the L3 and L2 edges, respectively. Similar to the pre-edge transitions in K-edge XAS, the intensities of these features reflect the total metal (n-1)d character in the unoccupied molecular orbitals, which means that increasing covalency will lead to lower L-edge intensities. This is a qualitative picture as the radial transition moment integral, which also contributes to the observed intensities, must also be considered.31 Beyond metal K- and L-edge XAS, ligand K-edge XAS is a very useful technique for insight into bonding in paramagnetic organometallic complexes.31 In this method, one utilizes the element specific nature of XAS to probe a specific atom bound to the metal in the organometallic complex of interest (e.g., N, S, Cl, etc.). While the absorption edge corresponds to the electric dipole allowed transition from the ligand 1s orbital to np, the pre-edge features will represent transitions from the ligand 1s orbital to unoccupied orbitals which can contain significant metal d character and therefore provide insight into metal-ligand covalency in an organometallic complex. Data obtained by this method is complementary to data obtained from metal XAS. At X-ray energies higher than the edge energy, the EXAFS region can provide information about the number, type and distance of neighboring atoms. At these energies the absorbing atom is ionized and emits a photoelectron of significant energy, with a De Broglie wavelength comparable to interatomic distances. The photoelectron propagates out of the absorbing atom as a wave and is backscattered by the electron density of surrounding atoms. The absorption of X-ray radiation is modulated by the interference between the outgoing and backscattered photoelectron waves.31,135,139 The photoelectron can be represented by the wavevector k: rffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi 2me ðE − E0 Þ k¼ (9) ħ2 where E is the energy of the electron, E0 is the binding energy, me is the electron mass and ħ is the reduced Planck’s constant. In the analysis of the EXAFS region, it is typical to define w(k) as the fractional modulation in the X-ray absorption coefficient: wðkÞ ¼

mðkÞ − m0 ðkÞ m0 ðkÞ

(10)

where m is the observed absorption coefficient and m0 is the absorption that would be observed in the absence of surrounding atoms. Since m0 cannot be directly measured it is approximated with the function: wðkÞ ¼

X Nf ðkÞ 2 2 e −2k s sin ½2kR + dðkފ 2 kR s

(11)

where N is number of scattering atoms of a given type, R is the distance between absorber and scatterer, f(k) is the amplitude function for the backscattering atom, s2 is the disorder parameter and d(k) is the phase shift for the absorber–backscatterer pair. The f(k) and d(k) parameters contain information about the identity of scattering atoms, but they depend weakly on the scatterer identity. As a result, atoms with similar numbers of electrons such as N/O or S/Cl cannot be differentiated whereas atoms with substantially different numbers of electrons such as O and S can be readily distinguished. From Eq. (11), it can be seen that amplitude is inversely proportional to R2 which means that the EXAFS analysis is limited to atoms in close proximity to the absorber. Hence, EXAFS oscillations are typically observed only for atoms within 5 A˚ of the absorber. Fourier transformation of EXAFS data transfers data from k-space to R-space which enables the visualization of radial distribution of electron density around the absorbing atom.135

1.06.4.3

Applications

One of the most common uses of XAS in the characterization of paramagnetic organometallic complexes is the determination of the oxidation state of the metal centers.140,142–146 For example, Meyer and coworkers used Cu K-edge XANES to assign the oxidation state of copper in a series of copper-NHC complexes (Fig. 23A). The high energy resolution fluorescence detected (HERFD) XAS spectra of the investigated complexes are shown in (Fig. 23B).147 The observed features are assigned using theoretical calculations. The spectrum of complex 1 shows two intense peaks at 8981.8 and 8984.8 eV, consistent with the completely filled d orbitals of Cu(I) metal center. The observed transitions were assigned to the electric dipole allowed 1s ! 4p transition. Due to the reduced symmetry of metal center the degeneracy of the Cu p orbitals is removed leading to the two observed transitions, where the first transition is a 1s ! 4py and second a 1s ! 4pz transition. It can be noted that the spectra of the other complexes are significantly different than the spectrum of complex 1, as they do not have intense electric dipole allowed transitions and the rising edge is shifted toward higher energies, which suggests a higher oxidation state of the metal center. From the spectrum of complex 2, a weak pre-edge transition at 8979.5 eV corresponds to the dipole forbidden 1s ! 3d transition to the only singly occupied orbital in the d9 Cu(II) center. In the spectrum of complex 3, both the pre-edge and the edge are shifted toward higher energy indicating the

Characterization Methods for Paramagnetic Organometallic Complexes

159

Fig. 23 (A) Structures of studied Cu complexes. (B) Cu Kb HERFD-XAS of studied complexes (10 K). Insert shows TD-DFT calculated pre-edge peaks for complexes 2 and 3 (dashed lines). Adapted with permission from Liu, Y.; Resch, S. G.; Klawitter, I.; Cutsail, G. E.; Demeshko, S.; Dechert, S.; Kühn, F. E.; DeBeer, S.; Meyer, F. Angew. Chem. Int. Ed. 2020, 59, 5696–5705. Copyright 2020 Wiley-VCH Verlag GmbH & Co. KGaA.

presence of a Cu(III) center in this complex. Calculations showed that the pre-edge features in the spectra of complexes 2 and 3 correspond to the transition from 1s to the 4dx2−y2 orbital. Additionally, the observed shift of the white line (intense absorptions on the rising edge) energies is also consistent with the increasing oxidation state of the Cu center in complexes 1–3. It is worth mentioning that the electronic structure of the formally d8 Cu(III) is a subject of active investigation and debate. It was shown that a majority of formal Cu(III) complexes exhibit inverted ligand fields in which unoccupied frontier molecular orbitals (FMOs) have predominantly ligand character.148 This means that in these complexes the metal center is better described as d10 Cu(I). Because of this it is necessary to determine metal d orbital character in unoccupied FMOs, experimentally (e.g., through Cu L2,3-edge XANES) and/or through theoretical calculations, prior to assigning a Cu(III) oxidation state. In addition to transition metal organometallic complexes, XAS can also be used for determining the oxidation states in organometallic lanthanide and actinide complexes.149–153 As previously mentioned, L3,2-edge XAS is primarily used for analysis of these compounds. In the work of Bart and coworkers, XAS was used to investigate the influence of different ligands on the

160

Characterization Methods for Paramagnetic Organometallic Complexes

electronic structure of a series of paramagnetic organometallic uranium compounds bearing trianionic redox active pyridine(diamine) ligands (MesPDIMe).154 The L-edge spectra of the investigated complexes, which is used for determining the effective nuclear charge on the uranium center, is shown in Fig. 24A, along with the spectrum of Cp UO2(MesPDIMe) (Cp ¼ 1,2,3,4,5-pentamethylcyclopentadienyl), a U(VI) complex which is used for comparison. Structures of the relevant complexes are shown on Fig. 24B. The inflection points of 3-Cp and 4-Cp (17172.1 and 17171.9 eV, respectively) are close to the inflection point of Cp UO2(MesPDIMe) (17171.9 eV), indicative of the same oxidation state of the metal center (in this case U(VI)). The spectra of 2-CpP (Cpp ¼ 1-(7,7-dimethylbenzyl)cyclopentadienyl) and 2-Cp exhibit a shift of the inflection point toward lower energy (0.3–0.9 eV), indicative of a lowering of the oxidation state of the uranium center. Even though the observed change in energy is smaller than previously observed by Meyer and coworkers (2.0 eV)155 for oxidation of U(V) to V(VI), based on other findings in this study complexes 2-CpP and 2-Cp are assigned as U(V) complexes. Comparing complexes with different ligands, it can be seen that the change from a phenyl(imido) ligand in 2-Cp to p-tolyl(imido) in 3-Cp leads to a change in the oxidation state of the metal center from U(V) to U(VI), which was also supported by NMR and UV-Vis experiments. Lastly, by comparing 3-Cp and its oxidized analog 4-Cp , the small change in inflection point energy observed was found to be suggestive of a ligand-based oxidation process, in accordance with other findings from this study. In a further example, Yang, Kozimor, Evans and coworkers combined L-edge XANES studies with theoretical calculations to evaluate the electronic structures of [LnII(C5H4SiMe3)3]− complexes (Ln ¼ Pr, Nd, Sm, Gd, Tb, Dy, Y, Ho, Er, Tm, Yb and Lu).156 Based on previous studies of these complexes, the electronic configurations of the lanthanide ions in [LnII(C5H4SiMe3)3]− (Ln ¼ La, Ce, Pr, Nd, Gd, Tb, Dy, Ho, Er, and Lu) anions were assigned as 4fn5d1 instead of the expected 4fn+15d0 configuration. Due to this unusual proposed ground state electronic structure, L-edge XANES was used to further probe these complexes as well as Sm(II), Tm(II), and Yb(II) analogs for which the electronic configurations are known to be 4f6, 4f13, and 4f14, respectively. Fig. 25 shows the XANES spectra of [SmII(C5H4SiMe3)3]− and SmII(C5Me5)2(THF)2 as well as the analogous complex, SmIII(C5H4SiMe3)3. The L3,2-edge peaks for [SmII(C5H4SiMe3)3]− are nearly identical to the L3,2-edge peaks for SmII(C5Me5)2(THF)2, as well as other previously characterized Sm(II) complexes. Comparison with the spectrum of the Sm(III) complex shows that L3,2-edge peaks are shifted toward higher energies, 7–8 eV, which agrees with previously reported XANES of Sm(II)/Sm(III) complexes.150,157–159 A similar situation is also present in the XANES spectra of Tm and Yb complexes. The XANES spectra of each of the studied complexes are shown on Fig. 26. When compared to the complexes of Yb, Sm and Tm, the shift of the L-edge between Ln(II) and Ln(III) complexes is much smaller (0.2–1 eV). This further supported the proposed 4fn5d1 electronic configuration in these complexes. The XANES spectra of the analogous Y and Lu complexes were also recorded. Reduction to the Y(II) complex places an electron in the 4d orbital, yielding a 4d1 electronic configuration. Similarly, the Lu(II) complex has a 4f145d1 electronic configuration, and the changes in the spectra of the Y and Lu complexes are similar to the changes observed for the Ln complexes with a 4fn5d1 electronic configuration. Taken together, these results imply that the reduction of the LnIII(C5H4SiMe3)3 complexes to the [LnII(C5H4SiMe3)3]− complexes resulted in the addition of an electron into a highly shielded 5d orbital, which results in the 4fn5d1 electronic configuration as opposed to the 4fn+15d0 configuration. For the analysis of organometallic complexes, ligand K-edge XAS can be used in combination with metal K-edge studies to provide further insight into electronic structure and bonding. As a representative example, Berry, DeBeer and coworkers used both S and Ni K-edge XAS to characterize electronic structure and bonding in a dinuclear Ni complex (Fig. 27A).160 The S K-edge XANES spectrum of complex 1 and its reduced version (1red) consist of an intense pre-edge peak followed by a broad rising K-edge (Fig. 27B). The pre-edge features arise from the excitation of an electron from the S 1s orbital to the unoccupied orbitals in the Ni 3d manifold which have Ni–S antibonding character. The observed shift of the pre-edge features for 1 and 1red suggest a Ni based reduction, which is further supported by similar energies of the rising edges. The complexes were further probed by Ni K-edge XANES. The pre-edge absorption can be observed around 8331 eV (Fig. 27C), and the difference in pre-edge energy for complexes 1 and 1red is small (0.1 eV) which is expected for isoleptic Ni(II)/Ni(III) complementary complexes with sulfur ligands.161 Even though the energy separation is small, the lower intensity after reduction indicates a Ni based reduction. Additional shoulders can be observed at 8338.3 eV for 1 and 8336.6 eV for 1red, each assigned as a 1s to 4p + LMCT (S to Ni) shakedown transition (Fig. 27D). This assignment is further supported by the spectrum of complex 4 which lacks S ligands and thus does not have this band. Furthermore, it can be observed that this band is lower in intensity and shifted toward lower energies upon reduction of 1, where the lower energy was attributed to the Ni based reduction, and the lower intensity attributed to a decreased covalency of the Ni–S bond in the reduced complex. As previously mentioned, XAS is a valuable tool for the determination of covalency in organometallic complexes.162–164 For example, Marin, Conardson, Clark and coworkers have used XAS to determine trends in covalency for d and f element complexes.165 The covalency of the M–Cl bond in a series of (C5Me5)2MCl2 (M ¼ Ti, 1; Zr, 2; Hf, 3; Th, 4; U, 5) complexes was studied using Cl K-edge XAS. In this experiment, the pre-edge features correspond to the ligand 1s electron excitation to the unoccupied orbitals of predominant metal character, which are mixed with ligand p orbitals. The integrated intensities of these pre-edge features serve as a direct measure of the ligand p character in these mixed orbitals, and thus a measurement of covalency. The Cl K-edge XAS spectrum of complexes 1–5 is shown in Fig. 28. The transition metal complexes (1–3) all have similar pre-edge features, with a shift toward higher energy when going from 3d to 4d to 5d.162 In contrast, the spectra of the actinide complexes are significantly different. The spectrum of the Th complex (4) shows no obvious pre-edge features, though second and fourth derivatives reveal a shoulder close to the rising edge. The uranium complex (5) shows three observable pre-edge features. The amount of Cl 3p character in the mixed orbitals for these complexes was determined from the corresponding pre-edge intensities via comparison to a Cs2CuCl4 standard which has 7.5% of Cl 3p character. The determined Cl 3p characters for complexes 1–3 are 25%, 24% and 22%, which are

Characterization Methods for Paramagnetic Organometallic Complexes

161

Fig. 24 (A) U L3,2-edge XANES spectra (77 K) and (B) structures of studied complex. Adapted with permission from Kiernicki, J. J.; Ferrier, M. G.; Lezama Pacheco, J. S.; La Pierre, H. S.; Stein, B. W.; Zeller, M.; Kozimor, S. A.; Bart, S. C. J. Am. Chem. Soc. 2016, 138, 13941–13951. Copyright 2020 American Chemical Society.

indistinguishable. While the pre-edge feature in the Th complex is not sufficiently resolved to enable determination of the peak intensity, the pre-edge features in for the U complex can be described by three peaks, where the two lower energy peaks are close to the position of Ti complex pre-edge peaks, and the third peak is higher in energy than first peak in Hf complex spectrum (which corresponds to transition into 5d orbital). This suggests that the third pre-edge peak for the U complex spectrum corresponds to a

162

Characterization Methods for Paramagnetic Organometallic Complexes

Fig. 25 Sm L3- and L2-edge XANES spectra of SmIII(C5H4SiMe3)3 (top, black), SmII(C5Me5)2(THF)2 (bottom, pink), and [K(2.2.2-cryptand)] [SmII(C5H4SiMe3)3] (bottom, black). Reprinted from Fieser, M. E.; Ferrier, M. G.; Su, J.; Batista, E.; Cary, S. K.; Engle, J. W.; Evans, W. J.; Lezama Pacheco, J. S.; Kozimor, S. A.; Olson, A. C.; Ryan, A. J.; Stein, B. W.; Wagner, G. L.; Woen, D. H.; Vitova, T.; Yang, P. Chem. Sci. 2017, 8, 6076–6091, with permission from the Royal Society of Chemistry.

transition to a 6d orbital, while the lower energy transitions potentially corresponded to the transitions into 5f orbitals. The determined Cl 3p character was 9%, which was more than two times lower than that observed for the transition metal complexes. Calculations confirmed that the pre-edge features in transition metal complexes correspond to the transitions to 3d (1), 4d (2) and 5d (3) orbitals. For complex 4, the 5f and 6d orbitals are close in energy which leads to a single pre-edge transition while for the U complex the first two pre-edge transitions were attributed to transitions into mainly 5f orbitals and the third transition corresponded to the transition into 6d orbitals. The calculated Cl 3p character follows the same order as determined experimentally, where covalency is lowered in the series Ti > Zr > Hf > U. This led to the conclusion that an increase in d orbital energies from 3d to 6d results in a decrease in covalency of M–Cl bonds due to the smaller interaction of orbitals. Beyond studies of well-defined paramagnetic organometallic complexes, XAS can also be used to characterize in situ formed species.29,166,167 Yamazoe, Tomotsu and coworkers used XAS to characterize the species in the titanium catalyzed syndiospecific styrene polymerization reaction.28 The reaction of (tBuC5H4)TiCl2(OAr) (Ar ¼ 2,6-diisopropylbenzene) with styrene in the presence of methylaluminoxane (MAO) as a cocatalyst was monitored by Ti K-edge XANES studies (Fig. 29). While the starting Ti(IV) complex exhibits two pre-edge features at 4966.5 and 4967.5 eV as well as shoulder on the edge at 4977.9 eV, the addition of MAO does not lead to the significant changes in the position of the pre-edge peaks and edge energies, suggesting that the oxidation state of the metal center remains Ti(IV). Addition of styrene, along with MAO resulted in the lowering of both the edge energy and the intensities of the pre-edge features, which was attributed to the reduction of the metal center to Ti(III). For further characterization, EXAFS showed that addition of both MAO and styrene resulted in significant changes in the EXAFS spectrum. Fitting of the EXAFS results enabled determination of the coordination numbers and distances of the surrounding ligand atoms. These results revealed that, after the reaction, chlorine is no longer coordinated to the metal center, and instead there is a new metal-carbon bond. Based on this result, the formed complex was hypothesized to contain both a cyclopentadienyl group and a phenoxide ligand, a conclusion which was further supported by theoretical calculations.

Characterization Methods for Paramagnetic Organometallic Complexes

163

Fig. 26 L-edge XANES spectra for LnIII(C5H4SiMe3)3 (black) and [K(2.2.2-cryptand)][LnII(C5H4SiMe3)3] (pink) for Ln ¼ Yb, Sm, Tm, Dy, Nd, Pr, Lu, Ho, Er, Tb and Gd. Spectra were collected at the Ln L3-edge except for Nd and Pr complexes, where spectra were collected at the L2-edge. Reprinted from Fieser, M. E.; Ferrier, M. G.; Su, J.; Batista, E.; Cary, S. K.; Engle, J. W.; Evans, W. J.; Lezama Pacheco, J. S.; Kozimor, S. A.; Olson, A. C.; Ryan, A. J.; Stein, B. W.; Wagner, G. L.; Woen, D. H.; Vitova, T.; Yang, P. Chem. Sci. 2017, 8, 6076–6091, with permission from the Royal Society of Chemistry.

In a further example of the application of XAS to in situ studies of organometallic complexes, Chirik, Delferro and coworkers investigated alcohol stability of bisphosphine cobalt precatalysts for asymmetric alkene hydrogenation (Fig. 30).168 Under catalytic conditions, (R,R)-(iPrDuPhos)Co(CH2SiMe3)2 ((R,R)-(iPrDuPhos) ¼ 1,2-bis((2R,5R)-2,5-di-i-propylphospholano)benzene), complex 1, is used as a precatalyst to achieve high hydrogenation yields in the presence of methanol. EXAFS was used to provide structural insights into the in situ formed cobalt species in catalysis. Starting with the precatalyst, the EXAFS analysis of 1 in THF supported a four-coordinate complex in solution with 2.1 Co −C and 2.1 Co −P scatterers, as expected for the bisphosphine dialkyl complex. Upon reaction with methanol, the number of Co–P scatterers is lowered to 1 indicating dissociation of a phosphine ligand. This reaction also leads to the formation of two new Co–O bonds, indicating ligation of methoxide and a solvent molecule to the Co complex during the reaction. The bisalkoxide complex with same chiral bisphosphine ligand, complex 2 was also studied. The EXAFS spectra of both the solid and THF solution of complex 2 showed the presence of two Co–P bonds and two Co–O bonds as expected. Upon reaction with methanol and similarly to complex 1, the number of Co–P scatterers is lowered to 0.5 while the number of Co–O scatterers is increased to 2.4, which is associated with the formation of the methanol ligated Co complex. These studies show that EXAFS can provide important insight into changes in the ligand environment of in situ generated species under catalytic conditions.

164

Characterization Methods for Paramagnetic Organometallic Complexes

Fig. 27 (A) Structure of studied complex 1. (B) S K-edge XANES spectra of 1 and 1red. (C) Ni K-edge spectra of 1, 1red, and 4 ([(Cp0 )2Ni]+) (20 K). (D) The first derivative of Ni K-edge spectra shown on (C). Adapted with permission from Yao, S. A.; Martin-Diaconescu, V.; Infante, I.; Lancaster, K. M.; Götz, A. W.; DeBeer, S.; Berry, J. F. J. Am. Chem. Soc. 2015, 137, 4993–5011. Copyright 2015 American Chemical Society.

Fig. 28 Cl K-edge XAS spectra of (C5Me5)2MCl2 (M ¼ Ti, Zr, Hf, Th, U) (room-temperature). Reprinted with permission from Kozimor, S. A.; Yang, P.; Batista, E. R.; Boland, K. S.; Burns, C. J.; Clark, D. L.; Conradson, S. D.; Martin, R. L.; Wilkerson, M. P.; Wolfsberg, L. E. J. Am. Chem. Soc. 2009, 131, 12125–12136. Copyright 2009 American Chemical Society.

Characterization Methods for Paramagnetic Organometallic Complexes

165

Fig. 29 Ti K-edge XANES spectra of (tBuC5H4)TiCl2(OAr) (3, Ar ¼ 2,6-iPr2C6H3) and after addition of MAO and styrene to the solution of 3 (in toluene, 25  C). Reprinted with permission from Nomura, K.; Izawa, I.; Yi, J.; Nakatani, N.; Aoki, H.; Harakawa, H.; Ina, T.; Mitsudome, T.; Tomotsu, N.; Yamazoe, S. Organometallics 2019, 38, 4497–4507. Copyright 2019 American Chemical Society.

1.06.5

Nuclear magnetic resonance

1.06.5.1

Introduction

In the NMR of paramagnetic compounds, the presence of unpaired electrons often results in large isotropic shifts (up to several hundred ppm) and/or broadened resonances, which often obscure any nuclear spin-spin coupling and make integration of signals difficult. The shifting and broadening in the NMR spectra of such systems arises from hyperfine interactions between unpaired electrons and the observed nucleus. Despite these complications, NMR of paramagnetic systems (often called paramagnetic NMR) offers special opportunities beyond those in NMR of diamagnetic compounds, as it can provide valuable information about spatial arrangement and bonding as well as delocalization of unpaired electron density. Other advantages to all NMR spectroscopy, such as insight into dynamics on the ms-ms timescale, are maintained despite the loss of information on coupling of nuclear spins. It is also applicable for monitoring reactions and studying in situ formed complexes. In addition, the “Evans method” described below can be used for determining the solution magnetic moment of paramagnetic complexes, which is useful for assigning spin of the ground state in mononuclear complexes.

1.06.5.2

Theory

Here, we skip the basics of NMR spectroscopy which can be found in literature169–173 and focus on interactions arising due to the presence of unpaired electrons in paramagnetic samples. The NMR active nuclei have a non-zero nuclear spin (nuclear spin angular momentum I) which gives rise to the nuclear magnetic moment (mI) analogous to the electron magnetic moment (Eq. 1)2,174: mI ¼ gI mN I ¼ ħgI I

(12)

where gI is the nuclear g factor, mN is the nuclear magneton, while gI is the nuclear gyromagnetic ratio. In an external magnetic field, the degeneracy of states with different MI values is removed (see Section 1.06.2.2). Interaction of an external magnetic field (B) with the nuclear magnetic moment can be described by nuclear Zeeman Hamiltonian:

Thus, the energy of any state is:

b Ze ¼ − mI B ¼ − ħgI BbI H

(13)

E ¼ − ħgI BMI

(14)

Due to the presence of unpaired electrons, beside nuclear Zeeman interaction, an additional electron Zeeman (Eq. 2) and hyperfine interaction (Eq. 5) must be considered. This causes further splitting of energy levels as shown on Fig. 31 for a simple model with

166 Characterization Methods for Paramagnetic Organometallic Complexes

Fig. 30 EXAFS spectra of studied Co complexes (black) and corresponding simulations (red) of (A) complex 1 in THF, (B) complex 2 in 9:1 THF:MeOH, (C) complex 2 in the solid state, (D) complex 2 in THF, (E) complex 2 in 9:1 THF:MeOH, (F) complex 3 in the solid state. Inserts are showing structures of studied complexes. Table below spectra are showing coordination numbers obtained from spectral fitting. aThe number was fixed in simulation. Reprinted with permission from Zhong, H.; Friedfeld, M. R.; Camacho-Bunquin, J.; Sohn, H.; Yang, C.; Delferro, M.; Chirik, P. J. Organometallics 2019, 38, 149–156. Copyright 2019 American Chemical Society.

Characterization Methods for Paramagnetic Organometallic Complexes

167

I ¼ 1/2 and S ¼ 1/2. As the NMR active nuclei are most commonly some distance away from the paramagnetic center, it is expected that hyperfine interaction is weaker than for EPR.174 Fig. 31 shows that considering only electron and nuclear Zeeman interactions results in two allowed NMR transitions (DMI ¼  1). Both transitions have the same energy (E ¼ ħgIB), which also corresponds to the transition energy for diamagnetic compounds (diamagnetic position, d on Fig. 31). The hyperfine interaction changes the relative energies of states involved in the transitions, which results in two transitions with different energies. These transitions are centered around the diamagnetic position and at A/2 distance from the center (A is hyperfine coupling constant). Even though it is expected to see two transitions in the spectrum, only one signal is observed in the NMR spectrum because the unpaired electrons relax on a timescale (10−14–10−8 s) that is several orders of magnitude shorter than the timescales of both nuclear relaxation and the NMR experiment (>10−3 s).2 If the population differences between states for both transitions are the same, that would result in just one signal at the diamagnetic position (due to the mentioned level averaging). However, considering the Boltzmann distribution, higher energy between states involved in the transitions results with a larger population difference, which means that one of the transitions will have higher intensity. That means that the signal in the spectrum is shifted from the diamagnetic position toward the higher intensity transition (Fig. 31). This shift from the diamagnetic position is known as a paramagnetic shift (or hyperfine shift, dpara T ), which is a temperature dependent measure of the hyperfine coupling constant A. Considering the interaction of an unpaired electron with the nucleus, the chemical shift measured in an NMR experiment can be represented as a sum of diamagnetic (ddia, not temperature dependent) and paramagnetic shifts (dpara T ): exp

para

dT ¼ ddia + dT

(15)

The hyperfine interaction can be divided into interactions coming from unpaired spin density located on the resonating nucleus, also known as Fermi contact interaction (or just contact interaction), and interactions from unpaired spin density outside the resonating nucleus, known as dipolar (or pseudocontact) interaction. This means that the paramagnetic shift can be represented as a

Fig. 31 The upper part of the figure shows energy level diagram for I ¼ 1/2 and S ¼ 1/2 system and NMR allowed transitions (red). The lower part of the figure shows resonances corresponding to allowed transitions (red) and observable resonance which is a consequence of level averaging (green).

168

Characterization Methods for Paramagnetic Organometallic Complexes

pc sum of the Fermi contact shift (dcon T ) and the pseudocontact shift (dT ). Unpaired electrons are not localized in space, rather they are delocalized in their corresponding molecular orbital (MO). Additionally, unpaired spin causes spin polarization. This means that paired elections in each filled MO are under the influence of unpaired electrons in other MOs in a such a way that electrons that are spin aligned with the unpaired electron will have a slight preference to occupy part of the MO closer to the unpaired electron. Similarly, the anti-aligned electrons have a slight preference to occupy part of the MO further away from the unpaired electron.175 This means that in the case of transition metals, unpaired electrons which are on metal d orbitals can induce delocalization of spin density at the NMR active nucleus. Delocalization of the unpaired electron spin density onto the nucleus causes the Fermi contact interaction (through-bond effect). Only unpaired spin density at the nucleus (in s orbitals, as only s orbitals have non-zero electron density at the nucleus) contributes to this interaction. In the simplified case, not considering ZFS and spin-orbit coupling, the Fermi contact shift can be represented by following equation175:

dcon T ¼

Acon gmB SðS + 1Þ 3gI kT ħ

(16)

where Acon is the contact hyperfine coupling constant, which is isotropic and related to the unpaired spin density at nucleus. Contact shifts are dominant for groups either directly bound to metal (or a few bonds away) or coupled to the metal strongly through a p-system. The pseudocontact shift arises due to the through-space spin-dipole coupling of unpaired electron density and nuclear spin. It is a consequence of anisotropic unpaired electron distribution. Again, in the simplified case, using the approximation that electron spin is localized at the metal center, the pseudocontact shift, for axial systems, can be represented as175: pc

dT ¼



m0 m2B SðS + 1Þ 1 2 3 cos 2 y − 1 g‖2 − g? 9kT r3 4p

(17)

where r is the electron-nuclear distance and y is the angle between r and the magnetic z-axis. The corresponding pseudocontact hyperfine coupling constant, Apc, is also a function of r and y, but unlike the contact shift, it can be anisotropic and provide insight into the location of the nucleus with respect to the magnetic axis.175–177 Therefore, in principle, it is possible to use the dpc T values to establish the direction of an atom if the magnetic axis is known, or to establish the direction of the magnetic axis if the locations of the nuclei are known. This method has been used widely in bioinorganic chemistry, but it is challenging in organometallic chemistry because there may be relatively few NMR-active nuclei and their locations may be dynamic. Estimating whether the contact or pseudocontact shift is the main contribution to the hyperfine shift is not trivial and depends on the precise nature of the studied system. If favorable orbital overlap leads to the significant unpaired spin delocalization to the nucleus of interest, then the contact shift is dominant, and if unpaired electron density is not close to the nucleus or the g tensor is particularly anisotropic (easy axis of magnetization), the pseudocontact shift is dominant.2 For example, in the case of lanthanide complexes the pseudocontact shift is typically dominant as large anisotropy of lanthanide ions leads to a strong dipolar interaction, and a small overlap of the contracted f-shell electrons with ligand orbitals results in a small contact shift.178 No matter which effect is dominant, the chemical shifts have a substantial temperature dependence, most often shifting away from the diamagnetic region (0–10 ppm) as the temperature is lowered. In many paramagnetic systems the proximity of unpaired electron density to the nuclear spin provides a facile mechanism for relaxation, known as paramagnetic relaxation enhancement.2 This leads to severe line broadening that influences resolution and hinders detection of all signals. The paramagnetic contributions to the T1 and T2 can be divided into contact and dipolar (pseudocontact) contributions, as well as the Curie spin contribution which arises from dipole-dipole interaction between the nucleus and the time averaged magnetic moment of electrons.4,179 Since most paramagnetic organometallic complexes are rapidly tumbling in solution, contact and dipolar contributions are the main source of paramagnetic relaxation enhancement. The nature of these interactions is the same as described before and more details about its influence on relaxation rates can be found in the literature.2 It is worth noting that line broadening due to the dipolar interaction drops off with r−6, unlike the pseudocontact shift which falls off with r−3 as previously shown. Additionally, both contributions are inversely proportional to the temperature, which means that increasing the temperature should lead to sharper resonances. Both contributions also directly depend on gI which means that nuclei with lower gI have sharper peaks (e.g., resonances in a 13C NMR spectrum are sharper than analogous resonances in a 1H NMR spectrum).4 In addition to the influences on its own nuclei, paramagnetic compounds in solution influence the chemical shifts of diamagnetic compounds in the same solution. This gives a convenient NMR based method for determining the effective magnetic moment of paramagnetic substances in solution, which is named the Evans method after its discoverer. In short, the application of an external magnetic field induces a bulk magnetization (total induced magnetic moment) which is proportional to the applied magnetic field, and a constant of that proportionality is the magnetic susceptibility (w).2 Evans showed that mass magnetic susceptibility can be related to the observed shift of resonances in a superconducting magnet using the following formula180: wm ¼

3Df w ðd0 − ds Þ + w0 + 0 2pfm m

(18)

where Df is the observed shift of a signal, f is the spectrometer frequency, m is the mass of the paramagnetic sample, w0 is the mass susceptibility of the solvent, d0 and ds are the densities of solution and the solvent, respectively. The last term in Eq. (18) is usually

Characterization Methods for Paramagnetic Organometallic Complexes

169

discarded for dilute solutions. Experimentally measured magnetic susceptibility is defined as a sum of paramagnetic (wpar) and diamagnetic (wdia) contributions. In order to determine paramagnetic susceptibility, the diamagnetic contribution has to be subtracted from the measured susceptibility. The diamagnetic contribution can be estimated by summing “Pascal’s constants” for diamagnetic susceptibility of every atom and bond in the molecule.181 Based on paramagnetic susceptibility, the effective magnetic moment can be calculated.182,183 qffiffiffiffiffiffiffiffiffiffiffiffiffi meff ¼ 8wpar T (19) The value for effective magnetic moment is typically compared to the “spin-only” magnetic moment expected for a given value of S pffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi spin − only ¼ 2 SðS + 1ÞmB ). Based on that comparison, the number of unpaired electrons can be inferred. It should be noted that (meff this solution value is determined at a temperature where low-lying excited states may be populated. Therefore, systems with magnetic coupling between paramagnetic centers give values that differ substantially from expectation. In these systems, other methods like SQUID magnetometry are recommended. Also, it should be remembered that small amounts of highly magnetic materials (iron and iron oxides for example) can give effects on the magnetization that are disproportionate to their mass, and are a frequent source of inaccuracy in Evans method interpretation.

1.06.5.3

Applications

Detailed information on the electronic structure of transition metal complexes typically requires EPR or MCD, which are advantageous as they can be used at low enough temperatures that the molecules are all in their ground states. NMR, on the other hand, is typically done at >150 K in solution which renders it less useful for electronic structure. However, NMR has major advantages: (a) measurement is rapid and routine, (b) the integration of signals is proportional to their concentrations, and thus it is useful for establishing purity. The collection parameters need to be adjusted: in addition to widening the sweep width, it is worthwhile to reduce the acquisition time so that more spectra can be averaged. Under these conditions, peaks from diamagnetic compounds are often saturated, so integrations should not be compared between paramagnetic and diamagnetic compounds. Also, when measuring concentrations of mixtures of paramagnetic compounds, it is worthwhile to add a capillary containing an integration standard (commonly Cp2Ni or Cp2Co, chosen to have no overlap with peaks from the mixture) because there may be species present with such broad peaks that they are not observable (“NMR silent”) and one can easily mistake a mixture for a pure compound. The following generalizations aside, paramagnetic NMR can be used in many cases to ascertain electronic structure details in combination with other techniques.99,184–186 For example, 1H NMR spectroscopy was used by Enders and coworkers to characterize quinolyl-functionalized Cr(III) complexes, which are precursors for highly active homogeneous Ziegler-Natta catalysts.187 In some cases, 1H NMR analysis of diamagnetic homogeneous Ziegler-Natta catalysts is limited due to the presence of high amounts of cocatalysts (like methylaluminoxane) as signals from the cocatalyst overlap with signals of catalysts. In this case, it is useful that the paramagnetic Cr(III) complexes have signals that are shifted away from the crowded diamagnetic region (0–10 ppm). Fig. 32A shows the 1H NMR spectrum of complex 1 in CDCl3 at room temperature along with the structure of this complex. The spectrum shows the presence of six paramagnetic peaks. Signal integration allowed the assignment of the peaks at 27.6 and −41.1 ppm to methyl groups (six H atoms). The difference in the position of these peaks is attributed to the fixed orientation of the Cp ring relative to the orbitals of the CrCl2 fragment, due to the coordination of the quinoline moiety to the metal center. Based on integration, the rest of the paramagnetic peaks are attributed to the H atoms on the quinoline moiety. A broader signal around −80 ppm is assigned to the H2 proton which is closest to the paramagnetic center resulting in the accentuated change in the chemical shift and peak

Fig. 32 1H NMR spectrum of (A) complex 1 and (B) complex 6 in CDCl3 (295 K, 200 MHz). Adapted with permission from Fernández, P.; Pritzkow, H.; Carbó, J. J.; Hofmann, P.; Enders, M. Organometallics 2007, 26, 4402–4412. Copyright 2007 American Chemical Society.

170

Characterization Methods for Paramagnetic Organometallic Complexes

broadening. The chemical shifts of the remaining proton resonances are in accordance with typical Fermi contact shift features where the sign of shift changes with the number of bonds between the paramagnetic center and the investigated nuclei.175,188 For an odd number of bonds a negative shift is expected, while for an even number of bonds the shift should be positive. Based on that, signals at 50.6 and −55.6 ppm were assigned to the H3 and H4, respectively. Similarly, signals at 15.3 and −16.4 ppm were assigned to H5 and H6, respectively. The signal from H7 could not be observed as it is most likely hidden in the diamagnetic region. For complex 6, where one of the methyl groups is replaced trimethylsilyl group, the spectrum showed similar features (Fig. 32B). It can be noted that the signal corresponding to the methyl groups at positions 9 and 10 is now split into two signals. The trimethylsilyl group is farther from the Cr center, and it is expected that the corresponding signals are in the diamagnetic region. The signal from the remaining methyl group could not be clearly observed at room temperature, but at higher temperatures the peaks became sharper and shifted toward the diamagnetic region and revealed a broad signal between 10 and 15 ppm assignable to the methyl group at position 11. Calculations based on Fermi contact shift yielded results that were in good correlation with experimental results, which indicated that this interaction is the main cause of paramagnetic shifts in these complexes. Paramagnetic NMR can be used to provide insight into spin delocalization as well as to provide information about the structure of investigated systems. Köhler, Sitzmann and coworkers used 1H and 13C NMR, along with EPR and magnetic studies, for studying binuclear Ni(II) complexes, [Cp000 Ni(m-Br)]2 (1; Cp000 ¼ 1,3,4-tBu3C5H2), [3CpNi(m-Br)]2 (2; 3Cp ¼ 1,3,4-iPr3C5H2), and [4CpNi(m-Br)]2 (3; 4Cp ¼ 2,3,4,5-iPr4C5H).189 Their resonances covered a wide range of chemical shifts in both 1H (550 ppm) and 13C NMR (2200 pm) NMR spectra. The spectra of investigated complexes are similar to the spectra of previously studied nickelocenes.190 In the 13C NMR spectra, resonances are more shifted toward higher frequencies, which is a consequence of direct delocalization of the spin density from the metal orbitals to the p orbitals of the Cp rings. Analysis of the influence of the substitution pattern on the splitting of the observed signals revealed that signal ordering in the 13C NMR spectra of trialkyl con con substituted complexes (1 and 2) has the following trend dcon T (C1) > dT (C3, 4) > dT (C2, 5), while the tetraalkylated complex (3) the signal trend is reversed (Fig. 33). Delocalization of the spin density also causes significant shifts in nuclei that are one (a), two (b) and three (g) bonds away from the Cp ring. The signals from the a nuclei have negative shifts and the ordering is determined, as in the previous case, from the analysis of the spin pattern on the Cp ring and supported by signal integration. The methyl groups at positions 3/4 in complex 2 and in complex 3 are diastereotopic, which results in the difference between b C and g H resonances. The spin transfer results in a positive spin at the b carbon and it depends on the dihedral angle y between the spin-carrying 2pz orbital of the carbon atom from the Cp ring and a–b bond. This means that the shift can report the dihedral angle y and the orientation of the substituent relative to the Cp ring. This analysis showed that the tBu groups in complex 1 and the iPr group at C1 position in complex 2 exhibited a dihedral angle close to 45 , as expected for alkyl groups bonded to a planar Cp ligand. The remaining iPr groups in complexes 2 and 3 have adjacent groups which constrain their rotation, leading to the deviation of the dihedral angle from 45 . It should be noted that the obtained orientations of iPr groups represent an equilibrium of orientations in solution. Based on the obtained data, signals coming from diastereotopic C nuclei were successfully assigned. The results were further used to assign proton resonances from diastereotopic methyl groups in complexes 2 and 3. Thus, this example showed how paramagnetic NMR can reveal the stereochemistry of paramagnetic organometallic complexes in solution. Paramagnetic NMR can also be successfully applied to lanthanide and actinide complexes.191–193 Enders and coworkers used 1H and 13C NMR for analysis of lanthanide (Tb, Dy, Ho, Er, Tm) complexes with cyclooctatetraene (COT) ligands.178 They used the structure-independent model in order to determine Fermi contact shift and pseudocontact shift contributions as well as magnetic anisotropy. Even though the analyzed complexes are paramagnetic, the observed resonances are not as broad as resonances for d-block metal complexes. Measurement at higher temperatures leads to even narrower peaks. For both studied ligands a maximum of 5 proton and 4 13C resonances could be observed, which points to the high symmetry of these complexes. The structures of the complexes with labeled H and C nuclei are shown on Fig. 34. In order to assign the observed resonances, the authors first considered nuclei that are separated by at least three bonds from the lanthanide ion and for which the Fermi contact interaction is small. For these nuclei, resonances are shifted mainly due to pseudocontact interactions. This allows, through geometric considerations, the estimation of the relative order of observed shifts. The chemical shift of the nuclei is proportional to the geometric factor (represents terms with r and y in Eq. 17), which is large for nuclei close to the metal center and close to the angle of 0 or 90 between the COT center, the metal ion and the examined nuclei. Considering the geometric factor, it is expected that the C3 resonance has a higher shift than C4. Also, inner nuclei H3i and H4i should have higher shifts than outer nuclei H3o and H4o. As the H3i and H3o nuclei are closer to the Ln ion it is expected that these signals are broader than H4i and H4o. Due to the proximity to the metal center, resonances of C1, C2 and H1 are expected to be significantly broadened, which prevents their detection in some cases. In the Dy complexes, the smaller chemical shifts indicate smaller magnetic anisotropy. The authors were able to assign all of the resonances based on the above-mentioned geometric considerations, except the C1 and C2 resonances which required a three-nucleus plot as described by Reuben.194 Calculations of the contributions to the resonance shifts revealed that considerable Fermi contact contributions are observable only for C atoms in the COT ring (C1 and C2). For other resonances the main contribution is a pseudocontact interaction. It was also determined that magnetic anisotropy is smallest for the Dy complex, in accordance with lower shifts (as the main interaction is a pseudocontact interaction), while the highest anisotropy was observed for the Tm complex. In a further example, Evans and coworkers used 29Si NMR on uranium complexes to determine if any trends exist between the chemical shift and the structure, ligand type, and oxidation state of the metal center.195 One of the reasons for establishing a correlation between oxidation state and chemical shifts is that the room temperature free ion magnetic moments of U4+ and U3+ are 3.62 and 3.58 mB, which is too close to be distinguished by room temperature magnetic susceptibility measurements. The 29Si NMR chemical shifts for various U complexes with oxidation states from +2 to +6 showed that aromatic ligands tend to give lower shifts than

Characterization Methods for Paramagnetic Organometallic Complexes

171

Fig. 33 1H NMR spectrum (top) of complex 1 in toluene-d8 (305 K) and the 13C NMR spectrum (bottom) 3 in toluene-d8 (375 K). Insert in the bottom spectrum shows expanded signals around −450 ppm. Signal labeled with S comes from the solvent. Reprinted with permission from Schär, M.; Saurenz, D.; Zimmer, F.; Schädlich, I.; Wolmershäuser, G.; Demeshko, S.; Meyer, F.; Sitzmann, H.; Heigl, O. M.; Köhler, F. H. Organometallics 2013, 32, 6298–6305. Copyright 2013 American Chemical Society.

Fig. 34 Structure of the [M(hdcCOT)2]− (left) and [M(odbCOT)2]− (right) anions with labeled nonequivalent positions. Reprinted with permission from Hiller, M.; Maier, M.; Wadepohl, H.; Enders, M. Organometallics 2016, 35, 1916–1922. Copyright 2016 American Chemical Society.

complexes with non-aromatic ligands. Additionally, in ligands where Si is bonded to N atoms 29Si NMR resonances are shifted more negative than in cases where Si is bonded to the C atom. It was also shown that 29Si resonances in cyclic ligands have more negative shifts. In terms of oxidation state, it was noted that, in most cases, the resonances of U4+ complexes are in range from 0 to −150 pm, while for U3+ complexes the resonances are mostly located between −120 and −250 ppm. The data suggest some useful trends but should not be viewed as definitive as they studied a limited number of complexes and 29Si NMR shifts are often variable.

172

Characterization Methods for Paramagnetic Organometallic Complexes

In addition to studying isolated compounds, NMR can be used for analyzing in situ formed species and tracking product formation in different chemical transformations. For example, Chirik and coworkers studied cobalt complexes bearing diisopropyl-substituted bis(phosphino)pyridine pincer ligand (iPrPNP), that can promote Suzuki–Miyaura cross coupling.196 These Co(I) complexes bearing pincer ligands were chosen as the flexibility of the ligand can accommodate both tetrahedral high spin (S ¼ 1) and planar low spin (S ¼ 0) Co(I) complexes. A series of aryloxide, alkoxide, and carboxylate complexes were synthesized and transmetalation with a neutral boron reagent was examined by 1H NMR. Addition of phenylboronic acid pinacol ester (PhBPin) to the (iPrPNP)CoOPh complex showed no changes in the NMR spectrum indicating no reaction. However, the addition of 2-benzofuranylBPin resulted in gradual formation of a new diamagnetic Co complex over the course of 24 h. Additional characterization revealed that the newly formed species is a diamagnetic, planar (iPrPNP)Co(2-benzofuranyl) complex, consistent with the observed changes in the 1H NMR spectrum of the reaction mixture. Similarly, the 1H NMR spectra for the reaction of 2-benzofuranylBPin with Co(I) alkoxide complexes showed rapid formation of a diamagnetic (iPrPNP)Co(2-benzofuranyl) complex. This demonstrates that the transmetalation step proceeds through rapid conversion of S ¼ 1 Co(I) alkoxides to S ¼ 0 Co(I) heteroarene products. Besides exploring reactivity, 1H NMR was also used to determine reaction rates for transmetalation of Co aryloxides and 2-benzofuranylBPin. Based on these results, the authors concluded that the relative rate constant increases with increasing electron donation from the aryl substituent. Additional kinetic experiments, performed by NMR, revealed that the overall reaction rate is second-order and first-order with respect to both cobalt and boron reagents. This example demonstrated the applicability of NMR for monitoring reaction progress and determination of reaction rates.

1.06.6

Conclusion

The development of base metal organometallic chemistry continues to increase the number of paramagnetic organometallic complexes, and their detailed characterization is important to define electronic structure, bonding and reactivity. This chapter has highlighted several spectroscopy methods that have proven to be especially useful for obtaining such insight into paramagnetic organometallics complexes. Specifically, EPR, MCD, XAS and NMR have been discussed in detail focusing on the fundamentals of each spectroscopic method as well as representative examples of the application of each method. The broader use of these methods will continue to expand our fundamental understanding of paramagnetic organometallic complexes and contribute to their applications in areas including catalysis.

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. 29. 30. 31. 32.

Albright, T. A. In Encyclopedia of Inorganic and Bioinorganic Chemistry; Scott, R. A., Ed.; John Wiley & Sons, Ltd: Chichester, UK, 2011. Pell, A. J.; Pintacuda, G.; Grey, C. P. Prog. Nucl. Magn. Reson. Spectrosc. 2019, 111, 1–271. Heise, H.; Köhler, F. H.; Herker, M.; Hiller, W. J. Am. Chem. Soc. 2002, 124, 10823–10832. Köhler, F. H. Encyclopedia of Magnetic Resonance; John Wiley & Sons, Ltd: Chichester, UK, 2011. Elsby, M. R.; Baker, R. T. Chem. Soc. Rev. 2020, 49, 8933–8987. Leznoff, D. B.; Mund, G. In Encyclopedia of Inorganic Chemistry; Scott, R. A., Ed.; John Wiley & Sons, Ltd: Chichester, UK, 2006. McKinney, R. J. Inorg. Chem. 1982, 21, 2051–2056. Alexander, M. D.; Harrington, P. C.; Van Heuvelen, A. J. Phys. Chem. 1971, 75, 3355–3359. MacAleese, L.; Maître, P. Mass Spectrom. Rev. 2007, 26, 583–605. Kahn, O. Molecular Magnetism; VCH: New York, NY, 1993. Heintz, R. A.; Koetzle, T. F.; Ostrander, R. L.; Rheingold, A. L.; Theopold, K. H.; Wu, P. Nature 1995, 378, 359–362. Ruiz, E.; Alvarez, S. Chem. Commun. 1998, (24), 2767–2768. Pariya, C.; Theopold, K. H. Curr. Sci. 2000, 78, 1345–1351. Lang, G. Le J. Phys. Colloq. 1971, 32, C1-822–C1-828. Testa-Anta, M.; Ramos-Docampo, M. A.; Comesaña-Hermo, M.; Rivas-Murias, B.; Salgueiriño, V. Nanoscale Adv. 2019, 1, 2086–2103. Swietlik, R.; Łapinski, A.; Fourmigué, M.; Yakushi, K. J. Raman Spectrosc. 2009, 40, 2092–2098. Stoian, S. A.; Smith, J. M.; Holland, P. L.; Münck, E.; Bominaar, E. L. Inorg. Chem. 2008, 47, 8687–8695. Rieger, A. L.; Rieger, P. H. Organometallics 2004, 23, 154–162. Van Doorslaer, S.; Murphy, D. M. In EPR Spectroscopy: Applications in Chemistry and Biology; Drescher, M., Jeschke, G., Eds.; Springer Berlin Heidelberg: Berlin, Heidelberg, 2011; pp 1–39. Goswami, M.; Chirila, A.; Rebreyend, C.; de Bruin, B. Top. Catal. 2015, 58, 719–750. Solomon, E. I.; Neidig, M. L.; Schenk, G. Comprehensive Coordination Chemistry II; Elsevier, 2003; vol. 2 pp 339–349. Braslavsky, S. E. Pure Appl. Chem. 2007, 79, 293–465. Wolford, N. J.; Radovic, A.; Neidig, M. L. Dalton Trans. 2021, 50, 416–428. Stephens, P. J. Annu. Rev. Phys. Chem. 1974, 25, 201–232. Carpenter, S. H.; Neidig, M. L. Isr. J. Chem. 2017, 57, 1106–1116. Oganesyan, V. S.; Thomson, A. J. J. Chem. Phys. 2000, 113, 5003. Kowalska, J.; DeBeer, S. Biochim. Biophys. Acta Mol. Cell Res. 1853, 2015, 1406–1415. Nomura, K.; Izawa, I.; Yi, J.; Nakatani, N.; Aoki, H.; Harakawa, H.; Ina, T.; Mitsudome, T.; Tomotsu, N.; Yamazoe, S. Organometallics 2019, 38, 4497–4507. Yi, J.; Nakatani, N.; Nomura, K. Dalton Trans. 2020, 49, 8008–8028. Campbell, W. J.; Brown, J. D.; Thatcher, J. W. Anal. Chem. 1966, 38, 416–439. Neese, F. In Comprehensive Inorganic Chemistry II; DeBeer, S., Reedijk, J., Poeppelmeier, K., Eds.; Elsevier, 2013; vol. 9; pp 427–439. Shankar, S.; Peters, M.; Steinborn, K.; Krahwinkel, B.; Sönnichsen, F. D.; Grote, D.; Sander, W.; Lohmiller, T.; Rüdiger, O.; Herges, R. Nat. Commun. 2018, 9, 4750.

Characterization Methods for Paramagnetic Organometallic Complexes

173

33. De, S.; Tewary, S.; Garnier, D.; Li, Y.; Gontard, G.; Lisnard, L.; Flambard, A.; Breher, F.; Boillot, M.-L.; Rajaraman, G.; Lescouëzec, R. Eur. J. Inorg. Chem. 2018, 2018, 414–428. 34. Celaje, J. A.; Pennington-Boggio, M. K.; Flaig, R. W.; Richmond, M. G.; Williams, T. J. Organometallics 2014, 33, 2019–2026. 35. Carsch, K. M.; DiMucci, I. M.; Iovan, D. A.; Li, A.; Zheng, S.-L.; Titus, C. J.; Lee, S. J.; Irwin, K. D.; Nordlund, D.; Lancaster, K. M.; Betley, T. A. Science (80-) 2019, 365, 1138–1143. 36. Schweiger, A.; Jeschke, G. Principles of Pulse Electron Paramagnetic Resonance; Oxford University Press, 2001. 37. Weil, J. A.; Bolton, J. R. Electron Paramagnetic Resonance; John Wiley & Sons, Inc.: Hoboken, NJ, USA, 2006; vol. 375. 38. Corvaja, C. In Electron Paramagnetic Resonance: A Practitioner’s Toolkit; Brustolon, M., Giamello, E., Eds.; John Wiley & Sons, Inc.: Hoboken, NJ, USA, 2008; pp 1–36. 39. Roessler, M. M.; Salvadori, E. Chem. Soc. Rev. 2018, 47, 2534–2553. 40. Gast, P.; Groenen, E. J. J. In EPR Spectroscopy: Fundamentals and Methods; Goldfarb, D., Stoll, S., Eds.; John Wiley & Sons, Ltd: Chichester, UK, 2018; pp 17–28. 41. Bennati, M. In EPR Spectroscopy: Fundamentals and Methods; Goldfarb, D., Stoll, S., Eds.; John Wiley & Sons Ltd: Chichester, UK, 2018; pp 81–94. 42. Hales, B. J. In Applications of Physical Methods to Inorganic and Bioinorganic Chemistry; Scott, R. A., Lukehart, C. M., Eds.; John Wiley & Sons Ltd: Chichester, UK, 2007; pp 39–54. 43. Boca, R. Coord. Chem. Rev. 2004, 248, 757–815. 44. Telser, J. In EPR Spectroscopy: Fundamentals and Methods; Goldfarb, D., Stoll, S., Eds.; John Wiley & Sons Ltd: Chichester, UK, 2018; pp 29–59. 45. van der Est, A. In EPR Spectroscopy: Fundamentals and Methods; Goldfab, D., Stoll, S., Eds.; John Wiley & Sons Ltd: Chichester, UK, 2018; pp 3–16. 46. Krzystek, J.; Zvyagin, S. A.; Ozarowski, A.; Trofimenko, S.; Telser, J. J. Magn. Reson. 2006, 178, 174–183. 47. Reiners, M.; Maekawa, M.; Baabe, D.; Zaretzke, M.-K.; Schweyen, P.; Daniliuc, C. G.; Freytag, M.; Raeder, J.; Hohenberger, J.; Sutter, J.; Meyer, K.; Walter, M. D. Dalton Trans. 2018, 47, 10517–10526. 48. Lichtenberg, C.; Viciu, L.; Adelhardt, M.; Sutter, J.; Meyer, K.; de Bruin, B.; Grützmacher, H. Angew. Chem. Int. Ed. 2015, 54, 5766–5771. 49. Hierlmeier, G.; Coburger, P.; Leest, N. P.; Bruin, B.; Wolf, R. Angew. Chem. Int. Ed. 2020, 59, 14148–14153. 50. Lohrey, T. D.; Rao, G.; Small, D. W.; Ouellette, E. T.; Bergman, R. G.; Britt, R. D.; Arnold, J. J. Am. Chem. Soc. 2020, 142, 13805–13813. 51. Lin, Q.; Diao, T. J. Am. Chem. Soc. 2019, 141, 17937–17948. 52. Townsend, T. M.; Bernskoetter, W. H.; Brudvig, G. W.; Hazari, N.; Lant, H. M. C.; Mercado, B. Q. Polyhedron 2020, 177, 114308. 53. Varela-Izquierdo, V.; López, J. A.; Bruin, B.; Tejel, C.; Ciriano, M. A. Chem. Eur. J. 2020, 26, 3270–3274. 54. Martinez, J. L.; Lutz, S. A.; Yang, H.; Xie, J.; Telser, J.; Hoffman, B. M.; Carta, V.; Pink, M.; Losovyj, Y.; Smith, J. M. Science 2020, 370, 356–359. 55. Bour, J. R.; Camasso, N. M.; Meucci, E. A.; Kampf, J. W.; Canty, A. J.; Sanford, M. S. J. Am. Chem. Soc. 2016, 138, 16105–16111. 56. Al-Afyouni, M. H.; Fillman, K. L.; Brennessel, W. W.; Neidig, M. L. J. Am. Chem. Soc. 2014, 136, 15457–15460. 57. Venkatasubbaiah, K.; Pakkirisamy, T.; Lalancette, R. A.; Jäkle, F. Dalton Trans. 2008, 0822, 4507. 58. Sakai, T.; Ohgo, Y.; Ikeue, T.; Takahashi, M.; Takeda, M.; Nakamura, M. J. Am. Chem. Soc. 2003, 125, 13028–13029. 59. Sakow, D.; Baabe, D.; Böker, B.; Burghaus, O.; Funk, M.; Kleeberg, C.; Menzel, D.; Pietzonka, C.; Bröring, M. Chem. Eur. J. 2014, 20, 2913–2924. 60. Muñoz, S. B.; Daifuku, S. L.; Brennessel, W. W.; Neidig, M. L. J. Am. Chem. Soc. 2016, 138, 7492–7495. 61. Aasa, R.; Vänngård, T. J. Magn. Reson. 1975, 19, 308–315. 62. Kostka, K. L.; Fox, B. G.; Hendrich, M. P.; Collins, T. J.; Rickard, C. E. F.; Wright, L. J.; Munck, E. J. Am. Chem. Soc. 1993, 115, 6746–6757. 63. Zultanski, S. L.; Fu, G. C. J. Am. Chem. Soc. 2013, 135, 624–627. 64. Jones, G. D.; Martin, J. L.; McFarland, C.; Allen, O. R.; Hall, R. E.; Haley, A. D.; Brandon, R. J.; Konovalova, T.; Desrochers, P. J.; Pulay, P.; Vicic, D. A. J. Am. Chem. Soc. 2006, 128, 13175–13183. 65. Phapale, V. B.; Buñuel, E.; García-Iglesias, M.; Cárdenas, D. J. Angew. Chem. Int. Ed. 2007, 46, 8790–8795. 66. Biswas, S.; Weix, D. J. J. Am. Chem. Soc. 2013, 135, 16192–16197. 67. Xu, H.; Zhao, C.; Qian, Q.; Deng, W.; Gong, H. Chem. Sci. 2013, 4, 4022. 68. Zheng, B.; Tang, F.; Luo, J.; Schultz, J. W.; Rath, N. P.; Mirica, L. M. J. Am. Chem. Soc. 2014, 136, 6499–6504. 69. Wu, J.; Nova, A.; Balcells, D.; Brudvig, G. W.; Dai, W.; Guard, L. M.; Hazari, N.; Lin, P.-H.; Pokhrel, R.; Takase, M. K. Chem. Eur. J. 2014, 20, 5327–5337. 70. Page, M. J.; Lu, W. Y.; Poulten, R. C.; Carter, E.; Algarra, A. G.; Kariuki, B. M.; Macgregor, S. A.; Mahon, M. F.; Cavell, K. J.; Murphy, D. M.; Whittlesey, M. K. Chem. Eur. J. 2013, 19, 2158–2167. 71. Zolnhofer, E. M.; Käß, M.; Khusniyarov, M. M.; Heinemann, F. W.; Maron, L.; van Gastel, M.; Bill, E.; Meyer, K. J. Am. Chem. Soc. 2014, 136, 15072–15078. 72. Zhang, Q.; Liu, Y.; Wang, T.; Zhang, X.; Long, C.; Wu, Y.-D.; Wang, M.-X. J. Am. Chem. Soc. 2018, 140, 5579–5587. 73. Mohadjer Beromi, M.; Nova, A.; Balcells, D.; Brasacchio, A. M.; Brudvig, G. W.; Guard, L. M.; Hazari, N.; Vinyard, D. J. J. Am. Chem. Soc. 2017, 139, 922–936. 74. Barth, E. L.; Davis, R. M.; Mohadjer Beromi, M.; Walden, A. G.; Balcells, D.; Brudvig, G. W.; Dardir, A. H.; Hazari, N.; Lant, H. M. C.; Mercado, B. Q.; Peczak, I. L. Organometallics 2019, 38, 3377–3387. 75. Hollmann, D.; Grabow, K.; Jiao, H.; Kessler, M.; Spannenberg, A.; Beweries, T.; Bentrup, U.; Brückner, A. Chem. Eur. J. 2013, 19, 13705–13713. 76. Xu, H.; Diccianni, J. B.; Katigbak, J.; Hu, C.; Zhang, Y.; Diao, T. J. Am. Chem. Soc. 2016, 138, 4779–4786. 77. Dzik, W. I.; Xu, X.; Zhang, X. P.; Reek, J. N. H.; de Bruin, B. J. Am. Chem. Soc. 2010, 132, 10891–10902. 78. Harris, C. F.; Bayless, M. B.; van Leest, N. P.; Bruch, Q. J.; Livesay, B. N.; Bacsa, J.; Hardcastle, K. I.; Shores, M. P.; de Bruin, B.; Soper, J. D. Inorg. Chem. 2017, 56, 12421–12435. 79. Tejel, C.; del Río, M. P.; Ciriano, M. A.; Reijerse, E. J.; Hartl, F.; Záliš, S.; Hetterscheid, D. G. H.; Tsichlis i Spithas, N.; de Bruin, B. Chem. Eur. J. 2009, (15), 11878–11889. 80. Jongbloed, L. S.; Vogt, N.; Sandleben, A.; de Bruin, B.; Klein, A.; van der Vlugt, J. I. Eur. J. Inorg. Chem. 2018, 2018, 2408–2418. 81. Butschke, B.; Fillman, K. L.; Bendikov, T.; Shimon, L. J. W.; Diskin-Posner, Y.; Leitus, G.; Gorelsky, S. I.; Neidig, M. L.; Milstein, D. Inorg. Chem. 2015, 54, 4909–4926. 82. Takaoka, A.; Peters, J. C. Inorg. Chem. 2012, 51, 16–18. 83. Zhang, G.; Wu, J.; Zheng, S.; Neary, M. C.; Mao, J.; Flores, M.; Trovitch, R. J.; Dub, P. A. J. Am. Chem. Soc. 2019, 141, 15230–15239. 84. Moehring, S. A.; Evans, W. J. Chem. Eur. J. 2020, 26, 1530–1534. 85. Huh, D. N.; Ziller, J. W.; Evans, W. J. Inorg. Chem. 2018, 57, 11809–11814. 86. Palumbo, C. T.; Halter, D. P.; Voora, V. K.; Chen, G. P.; Chan, A. K.; Fieser, M. E.; Ziller, J. W.; Hieringer, W.; Furche, F.; Meyer, K.; Evans, W. J. Inorg. Chem. 2018, 57, 2823–2833. 87. Wolford, N. J.; Yu, X.; Bart, S. C.; Autschbach, J.; Neidig, M. L. Dalton Trans. 2020, 49, 14401–14410. 88. Hümmer, J.; Heinemann, F. W.; Meyer, K. Inorg. Chem. 2017, 56, 3201–3206. 89. Langeslay, R. R.; Chen, G. P.; Windorff, C. J.; Chan, A. K.; Ziller, J. W.; Furche, F.; Evans, W. J. J. Am. Chem. Soc. 2017, 139, 3387–3398. 90. Boreen, M. A.; Rao, G.; Villarreal, D. G.; Watt, F. A.; Britt, R. D.; Hohloch, S.; Arnold, J. Chem. Commun. 2020, 56, 4535–4538. 91. Moehring, S. A.; Evans, W. J. Organometallics 2020, 39, 1187–1194. 92. Corbey, J. F.; Woen, D. H.; Palumbo, C. T.; Fieser, M. E.; Ziller, J. W.; Furche, F.; Evans, W. J. Organometallics 2015, 34, 3909–3921. 93. Andres, H.; Bominaar, E. L.; Smith, J. M.; Eckert, N. A.; Holland, P. L.; Münck, E. J. Am. Chem. Soc. 2002, 124, 3012–3025. 94. Savitsky, A.; Möbius, K. Photosynth. Res. 2009, 102, 311–333. 95. Krzystek, J.; Ozarowski, A.; Telser, J. Coord. Chem. Rev. 2006, 250, 2308–2324.

174

Characterization Methods for Paramagnetic Organometallic Complexes

96. Bucinsky, L.; Breza, M.; Lee, W.-T.; Hickey, A. K.; Dickie, D. A.; Nieto, I.; DeGayner, J. A.; Harris, T. D.; Meyer, K.; Krzystek, J.; Ozarowski, A.; Nehrkorn, J.; Schnegg, A.; Holldack, K.; Herber, R. H.; Telser, J.; Smith, J. M. Inorg. Chem. 2017, 56, 4751–4768. 97. Tran, B. L.; Singhal, M.; Park, H.; Lam, O. P.; Pink, M.; Krzystek, J.; Ozarowski, A.; Telser, J.; Meyer, K.; Mindiola, D. J. Angew. Chem. Int. Ed. 2010, 49, 9871–9875. 98. Jackson, T. A.; Krzystek, J.; Ozarowski, A.; Wijeratne, G. B.; Wicker, B. F.; Mindiola, D. J.; Telser, J. Organometallics 2012, 31, 8265–8274. 99. Krzystek, J.; Kohl, G.; Hansen, H.-B.; Enders, M.; Telser, J. Organometallics 2019, 38, 2179–2188. 100. Stalzer, M. M.; Telser, J.; Krzystek, J.; Motta, A.; Delferro, M.; Marks, T. J. Organometallics 2016, 35, 2683–2688. 101. Pascual-Álvarez, A.; Vallejo, J.; Pardo, E.; Julve, M.; Lloret, F.; Krzystek, J.; Armentano, D.; Wernsdorfer, W.; Cano, J. Chem. Eur. J. 2015, 21, 17299–17307. 102. Krzystek, J.; Telser, J.; Hoffman, B. M.; Brunel, L.-C.; Licoccia, S. J. Am. Chem. Soc. 2001, 123, 7890–7897. 103. Krzystek, J.; Telser, J.; Pardi, L. A.; Goldberg, D. P.; Hoffman, B. M.; Brunel, L.-C. Inorg. Chem. 1999, 38, 6121–6129. 104. Stoll, S. In EPR Spectroscopy: Fundamentals and Methods; Goldfab, D., Stoll, S., Eds.; John Wiley & Sons Ltd: Chichester, UK, 2018; pp 215–232. 105. Bowman, M. K. In Electron Paramagnetic Resonance: A Practitioner’s Toolkit; Brustolon, M., Giamello, E., Eds.; John Wiley & Sons, Inc.: Hoboken, NJ, USA, 2008; pp 159–194. 106. Dikanov, S. A.; Tsvetkov, Y. Electron Spin Echo Envelope Modulation (ESEEM) Spectroscopy; ESEEM Spectroscopy; CRC Press, 1992. 107. Mims, W. B. Phys. Rev. B 1972, 5, 2409–2419. 108. Höfer, P.; Grupp, A.; Nebenführ, H.; Mehring, M. Chem. Phys. Lett. 1986, 132, 279–282. 109. Königsmann, M.; Donati, N.; Stein, D.; Schönberg, H.; Harmer, J.; Sreekanth, A.; Grützmacher, H. Angew. Chem. Int. Ed. 2007, 46, 3567–3570. 110. Citek, C.; Oyala, P. H.; Peters, J. C. J. Am. Chem. Soc. 2019, 141, 15211–15221. 111. Lichtenberg, C.; Garcia Rubio, I.; Viciu, L.; Adelhardt, M.; Meyer, K.; Jeschke, G.; Grützmacher, H. Angew. Chem. Int. Ed. 2015, 54, 13012–13017. 112. Allouche, F.; Klose, D.; Gordon, C. P.; Ashuiev, A.; Wörle, M.; Kalendra, V.; Mougel, V.; Copéret, C.; Jeschke, G. Angew. Chem. Int. Ed. 2018, 57, 14533–14537. 113. Formanuik, A.; Ariciu, A. M.; Ortu, F.; Beekmeyer, R.; Kerridge, A.; Tuna, F.; McInnes, E. J. L.; Mills, D. P. Nat. Chem. 2017, 9, 578–583. 114. Mack, J.; Stillman, M. J. In Encyclopedia of Inorganic and Bioinorganic Chemistry; Scott, R. A., Ed.; John Wiley & Sons, Ltd: Chichester, UK, 2011. 115. Neese, F.; Solomon, E. I. Inorg. Chem. 1999, 38, 1847–1865. 116. Pavel, E. G.; Kitajima, N.; Solomon, E. I. J. Am. Chem. Soc. 1998, 120, 3949–3962. 117. O’Brien, H. M.; Manzotti, M.; Abrams, R. D.; Elorriaga, D.; Sparkes, H. A.; Davis, S. A.; Bedford, R. B. Nat. Catal. 2018, 1, 429–437. 118. Loup, J.; Zell, D.; Oliveira, J. C. A.; Keil, H.; Stalke, D.; Ackermann, L. Angew. Chem. Int. Ed. 2017, 56, 14197–14201. 119. Guisán-Ceinos, M.; Tato, F.; Buñuel, E.; Calle, P.; Cárdenas, D. J. Chem. Sci. 2013, 4, 1098. 120. Deng, Y.; Wei, X.; Wang, X.; Sun, Y.; Noël, T. Chem. Eur. J. 2019, 25, 14532–14535. 121. Li, B.-J.; Xu, L.; Wu, Z.-H.; Guan, B.-T.; Sun, C.-L.; Wang, B.-Q.; Shi, Z.-J. J. Am. Chem. Soc. 2009, 131, 14656–14657. 122. Fillman, K. L.; Przyojski, J. A.; Al-Afyouni, M. H.; Tonzetich, Z. J.; Neidig, M. L. Chem. Sci. 2015, 6, 1178–1188. 123. Muñoz, S. B.; Fleischauer, V. E.; Brennessel, W. W.; Neidig, M. L. Organometallics 2018, 37, 3093–3101. 124. Baker, T. M.; Mako, T. L.; Vasilopoulos, A.; Li, B.; Byers, J. A.; Neidig, M. L. Organometallics 2016, 35, 3692–3700. 125. Iannuzzi, T. E.; Gao, Y.; Baker, T. M.; Deng, L.; Neidig, M. L. Dalton Trans. 2017, 46, 13290–13299. 126. Krzystek, J.; Zvyagin, S. A.; Ozarowski, A.; Fiedler, A. T.; Brunold, T. C.; Telser, J. J. Am. Chem. Soc. 2004, 126, 2148–2155. 127. Kato, H.; Akimoto, K. J. Am. Chem. Soc. 1974, 96, 1351–1357. 128. Fleischauer, V. E.; Ganguly, G.; Woen, D. H.; Wolford, N. J.; Evans, W. J.; Autschbach, J.; Neidig, M. L. Organometallics 2019, 38, 3124–3131. 129. Wolford, N. J.; Sergentu, D.; Brennessel, W. W.; Autschbach, J.; Neidig, M. L. Angew. Chem. Int. Ed. 2019, 58, 10266–10270. 130. Daifuku, S. L.; Al-Afyouni, M. H.; Snyder, B. E. R.; Kneebone, J. L.; Neidig, M. L. J. Am. Chem. Soc. 2014, 136, 9132–9143. 131. Daifuku, S. L.; Kneebone, J. L.; Snyder, B. E. R.; Neidig, M. L. J. Am. Chem. Soc. 2015, 137, 11432–11444. 132. Camasso, N. M.; Canty, A. J.; Ariafard, A.; Sanford, M. S. Organometallics 2017, 36, 4382–4393. 133. Watson, M. B.; Rath, N. P.; Mirica, L. M. J. Am. Chem. Soc. 2017, 139, 35–38. 134. Shin, J.; Gwon, S.; Kim, S.; Lee, J.; Park, K. J. Am. Chem. Soc. 2020, 142, 4173–4183. 135. Penner-Hahn, J. E. In Comprehensive Coordination Chemistry II; McCleverty, J. A., Meyer, T. J., Eds.; Elsevier, 2003; vol. 2; pp 159–186. 136. Yano, J.; Yachandra, V. K. Photosynth. Res. 2009, 102, 241–254. 137. Baker, M. L.; Mara, M. W.; Yan, J. J.; Hodgson, K. O.; Hedman, B.; Solomon, E. I. Coord. Chem. Rev. 2017, 345, 182–208. 138. Bencze, K. Z.; Kondapalli, K. C.; Stemmler, T. L. Encyclopedia of Inorganic and Bioinorganic Chemistry; John Wiley & Sons, Ltd: Chichester, UK, 2011. 139. Smith, J. W.; Saykally, R. J. Chem. Rev. 2017, 117, 13909–13934. 140. Nelson, R. C.; Miller, J. T. Catal. Sci. Technol. 2012, 2, 461–470. 141. Sarangi, R. Coord. Chem. Rev. 2013, 257, 459–472. 142. Darmon, J. M.; Stieber, S. C. E.; Sylvester, K. T.; Fernández, I.; Lobkovsky, E.; Semproni, S. P.; Bill, E.; Wieghardt, K.; DeBeer, S.; Chirik, P. J. J. Am. Chem. Soc. 2012, 134, 17125–17137. 143. Yu, R. P.; Darmon, J. M.; Milsmann, C.; Margulieux, G. W.; Stieber, S. C. E.; DeBeer, S.; Chirik, P. J. J. Am. Chem. Soc. 2013, 135, 13168–13184. 144. Danopoulos, A. A.; Braunstein, P.; Monakhov, K. Y.; van Leusen, J.; Kögerler, P.; Clémancey, M.; Latour, J.-M.; Benayad, A.; Tromp, M.; Rezabal, E.; Frison, G. Dalton Trans. 2017, 46, 1163–1171. 145. MacLeod, K. C.; DiMucci, I. M.; Zovinka, E. P.; McWilliams, S. F.; Mercado, B. Q.; Lancaster, K. M.; Holland, P. L. Organometallics 2019, 38, 4224–4232. 146. Yogendra, S.; Weyhermüller, T.; Hahn, A. W.; Debeer, S. Inorg. Chem. 2019, 58, 9358–9367. 147. Liu, Y.; Resch, S. G.; Klawitter, I.; Cutsail, G. E.; Demeshko, S.; Dechert, S.; Kühn, F. E.; DeBeer, S.; Meyer, F. Angew. Chem. Int. Ed. 2020, 59, 5696–5705. 148. DiMucci, I. M.; Lukens, J. T.; Chatterjee, S.; Carsch, K. M.; Titus, C. J.; Lee, S. J.; Nordlund, D.; Betley, T. A.; MacMillan, S. N.; Lancaster, K. M. J. Am. Chem. Soc. 2019, 141, 18508–18520. 149. Nocton, G.; Booth, C. H.; Maron, L.; Andersen, R. A. Organometallics 2013, 32, 5305–5312. 150. Harder, S.; Naglav, D.; Ruspic, C.; Wickleder, C.; Adlung, M.; Hermes, W.; Eul, M.; Pöttgen, R.; Rego, D. B.; Poineau, F.; Czerwinski, K. R.; Herber, R. H.; Nowik, I. Chem. Eur. J. 2013, 19, 12272–12280. 151. Pattenaude, S. A.; Mullane, K. C.; Schelter, E. J.; Ferrier, M. G.; Stein, B. W.; Bone, S. E.; Lezama Pacheco, J. S.; Kozimor, S. A.; Fanwick, P. E.; Zeller, M.; Bart, S. C. Inorg. Chem. 2018, 57, 6530–6539. 152. Minasian, S. G.; Krinsky, J. L.; Rinehart, J. D.; Copping, R.; Tyliszczak, T.; Janousch, M.; Shuh, D. K.; Arnold, J. J. Am. Chem. Soc. 2009, 131, 13767–13783. 153. Meihaus, K. R.; Minasian, S. G.; Lukens, W. W.; Kozimor, S. A.; Shuh, D. K.; Tyliszczak, T.; Long, J. R. J. Am. Chem. Soc. 2014, 136, 6056–6068. 154. Kiernicki, J. J.; Ferrier, M. G.; Lezama Pacheco, J. S.; La Pierre, H. S.; Stein, B. W.; Zeller, M.; Kozimor, S. A.; Bart, S. C. J. Am. Chem. Soc. 2016, 138, 13941–13951. 155. Kosog, B.; La Pierre, H. S.; Denecke, M. A.; Heinemann, F. W.; Meyer, K. Inorg. Chem. 2012, 51, 7940–7944. 156. Fieser, M. E.; Ferrier, M. G.; Su, J.; Batista, E.; Cary, S. K.; Engle, J. W.; Evans, W. J.; Lezama Pacheco, J. S.; Kozimor, S. A.; Olson, A. C.; Ryan, A. J.; Stein, B. W.; Wagner, G. L.; Woen, D. H.; Vitova, T.; Yang, P. Chem. Sci. 2017, 8, 6076–6091. 157. Agondanou, J.-H.; Spyroulias, G. A.; Purans, J.; Tsikalas, G.; Souleau, C.; Coutsolelos, A. G.; Bénazeth, S. Inorg. Chem. 2001, 40, 6088–6096. 158. Menushenkov, A. P.; Chernikov, R. V.; Sidorov, V. V.; Klementiev, K. V.; Alekseev, P. A.; Rybina, A. V. JETP Lett. 2006, 84, 119–123. 159. Beaurepaire, E.; Kappler, J. P.; Krill, G. Phys. Rev. B 1990, 41, 6768–6776. 160. Yao, S. A.; Martin-Diaconescu, V.; Infante, I.; Lancaster, K. M.; Götz, A. W.; DeBeer, S.; Berry, J. F. J. Am. Chem. Soc. 2015, 137, 4993–5011. 161. Colpas, G. J.; Maroney, M. J.; Bagyinka, C.; Kumar, M.; Willis, W. S.; Suib, S. L.; Mascharak, P. K.; Baidya, N. Inorg. Chem. 1991, 30, 920–928.

Characterization Methods for Paramagnetic Organometallic Complexes

175

162. Kozimor, S. A.; Yang, P.; Batista, E. R.; Boland, K. S.; Burns, C. J.; Christensen, C. N.; Clark, D. L.; Conradson, S. D.; Hay, P. J.; Lezama, J. S.; Martin, R. L.; Schwarz, D. E.; Wilkerson, M. P.; Wolfsberg, L. E. Inorg. Chem. 2008, 47, 5365–5371. 163. Minasian, S. G.; Keith, J. M.; Batista, E. R.; Boland, K. S.; Clark, D. L.; Kozimor, S. A.; Martin, R. L.; Shuh, D. K.; Tyliszczak, T. Chem. Sci. 2014, 5, 351–359. 164. Smiles, D. E.; Batista, E. R.; Booth, C. H.; Clark, D. L.; Keith, J. M.; Kozimor, S. A.; Martin, R. L.; Minasian, S. G.; Shuh, D. K.; Stieber, S. C. E.; Tyliszczak, T. Chem. Sci. 2020, 11, 2796–2809. 165. Kozimor, S. A.; Yang, P.; Batista, E. R.; Boland, K. S.; Burns, C. J.; Clark, D. L.; Conradson, S. D.; Martin, R. L.; Wilkerson, M. P.; Wolfsberg, L. E. J. Am. Chem. Soc. 2009, 131, 12125–12136. 166. Bartlett, S. A.; Besley, N. A.; Dent, A. J.; Diaz-Moreno, S.; Evans, J.; Hamilton, M. L.; Hanson-Heine, M. W. D.; Horvath, R.; Manici, V.; Sun, X.-Z.; Towrie, M.; Wu, L.; Zhang, X.; George, M. W. J. Am. Chem. Soc. 2019, 141, 11471–11480. 167. van Bokhoven, J. A.; Lamberti, C. In Nanotechnology in Catalysis: Applications in the Chemical Industry, Energy Development, and Environment Protection; Van de Voorde, M., Sels, B., Eds.; Wiley-VCH Verlag GmbH & Co. KGaA: Weinheim, Germany, 2017; pp 1029–1054. 168. Zhong, H.; Friedfeld, M. R.; Camacho-Bunquin, J.; Sohn, H.; Yang, C.; Delferro, M.; Chirik, P. J. Organometallics 2019, 38, 149–156. 169. Jacobsen, N. E. NMR Spectroscopy Explained: Simplified Theory, Applications and Examples for Organic Chemistry and Structural Biology; John Wiley & Sons, Inc.: Hoboken, NJ, USA, 2007. 170. Günther, H. NMR Spectroscopy: Basic Principles, Concepts, and Applications in Chemistry; Wiley-VCH Verlag GmbH & Co. KGaA, 2013. 171. Keeler, J. Understanding NMR Spectroscopy; John Wiley & Sons, Ltd, 2010. 172. Levitt, M. H. Spin Dynamics: Basics of Nuclear Magnetic Resonance; John Wiley & Sons, Ltd, 2008. 173. Lambert, J. B.; Mazzola, E. P.; Ridge, C. D. Nuclear Magnetic Resonance Spectroscopy: An Introduction to Principles, Applications, and Experimental Methods; John Wiley & Sons, Ltd, 2019. 174. Jesson, J. P. In NMR of Paramagnetic Molecules; La Mar, G. N., Horrocks, W. D., Holm, R. H., Eds.; Elsevier, 1973; pp 1–52. 175. Solution NMR of Paramagnetic Molecules. Bertini, I., Luchinat, C., Parigi, G., Eds.; In Current Methods in Inorganic Chemistry; Elsevier, 2001; vol. 2; pp 29–73. 176. Bertini, I.; Luchinat, C.; Parigi, G. Concepts Magn. Reson. 2002, 14, 259–286. 177. Gochin, M.; Roder, H. Protein Sci. 2008, 4, 296–305. 178. Hiller, M.; Maier, M.; Wadepohl, H.; Enders, M. Organometallics 2016, 35, 1916–1922. 179. Marius Clore, G.; Iwahara, J. Chem. Rev. 2009, 109, 4108–4139. 180. Evans, D. F. J. Chem. Soc. 1959, 2003. 181. Bain, G. A.; Berry, J. F. J. Chem. Educ. 2008, 85, 532. 182. Hoppe, J. I. J. Chem. Educ. 1972, 49, 505. 183. Piguet, C. J. Chem. Educ. 1997, 74, 815. 184. Soshnikov, I. E.; Semikolenova, N. V.; Bryliakov, K. P.; Zakharov, V. A.; Sun, W.-H.; Talsi, E. P. Organometallics 2015, 34, 3222–3227. 185. Gordon, C. P.; Yamamoto, K.; Liao, W.-C.; Allouche, F.; Andersen, R. A.; Copéret, C.; Raynaud, C.; Eisenstein, O. ACS Cent. Sci. 2017, 3, 759–768. 186. Gordon, C. P.; Raynaud, C.; Andersen, R. A.; Copéret, C.; Eisenstein, O. Acc. Chem. Res. 2019, 52, 2278–2289. 187. Fernández, P.; Pritzkow, H.; Carbó, J. J.; Hofmann, P.; Enders, M. Organometallics 2007, 26, 4402–4412. 188. Wilkens, S. J.; Xia, B.; Weinhold, F.; Markley, J. L.; Westler, W. M. J. Am. Chem. Soc. 1998, 120, 4806–4814. 189. Schär, M.; Saurenz, D.; Zimmer, F.; Schädlich, I.; Wolmershäuser, G.; Demeshko, S.; Meyer, F.; Sitzmann, H.; Heigl, O. M.; Köhler, F. H. Organometallics 2013, 32, 6298–6305. 190. Blümel, J.; Hofmann, P.; Köhler, F. H. Magn. Reson. Chem. 1993, 31, 2–6. 191. Huang, W.; Upton, B. M.; Khan, S. I.; Diaconescu, P. L. Organometallics 2013, 32, 1379–1386. 192. Clark, C. L.; Lockhart, J. J.; Fanwick, P. E.; Bart, S. C. Chem. Commun. 2015, 51, 14084–14087. 193. Tsoureas, N.; Mansikkamäki, A.; Layfield, R. A. Chem. Commun. 2020, 56, 944–947. 194. Reuben, J. J. Magn. Reson. 1982, 50, 233–236. 195. Windorff, C. J.; Evans, W. J. Organometallics 2014, 33, 3786–3791. 196. Neely, J. M.; Bezdek, M. J.; Chirik, P. J. ACS Cent. Sci. 2016, 2, 935–942.

1.07

Computational Methods in Organometallic Chemistry

S Chantal E Stieber, Department of Chemistry & Biochemistry, California State Polytechnic University, Pomona, CA, United States Copyright © 2022 Elsevier Ltd. All rights reserved. This Contribution is published under the Creative Commons Attribution-Noncommercial 4.0 International (CC BY-NC 4.0) licensing conditions.

1.07.1 Introduction 1.07.2 Density functional theory (DFT) and time-dependent density functional theory (TD-DFT) 1.07.2.1 Functional 1.07.2.2 Basis set 1.07.2.3 Additional considerations 1.07.2.4 Beginning calculations 1.07.2.4.1 Geometries and geometry optimization 1.07.2.4.2 Single point calculation 1.07.2.4.3 Software and common programs 1.07.2.5 Closed Shell systems and restricted Kohn-Sham (RKS) 1.07.2.6 Open Shell systems and unrestricted Kohn-Sham (UKS) 1.07.2.7 Broken symmetry calculations 1.07.2.7.1 Case study 1: Broken symmetry calculations of (iPrPDI)FeN2 1.07.2.8 Determining the correct solution 1.07.2.9 F-elements 1.07.2.10 Limitations of DFT 1.07.3 Ab initio methods 1.07.3.1 Case study 2: Multiconfigurational calculations of an iron nitrosyl complex 1.07.4 Configuration interaction/multiplet calculations 1.07.4.1 Simple configuration interaction model for spectroscopy 1.07.4.2 Charge transfer model for spectroscopy 1.07.4.2.1 Case study 3: Multiplet theory for quantifying lanthanide covalency 1.07.5 Experimental applications 1.07.5.1 Benchmarking computations with experimental data 1.07.5.2 UV–visible spectroscopy 1.07.5.2.1 Case study 4: Calculated photodynamics and UV–visible spectroscopy in a ruthenium nitrosyl complex 1.07.5.3 Infrared spectroscopy 1.07.5.3.1 Case study 5: DFT calculations of infrared spectra for iron nitrosyl complexes 1.07.5.4 Nuclear magnetic resonance spectroscopy 1.07.5.4.1 Case study 6: Computationally predicting olefin metathesis intermediates with 13C NMR spectroscopy 1.07.5.5 Electron paramagnetic resonance spectroscopy 1.07.5.5.1 Computational methods for electron paramagnetic resonance spectroscopy 1.07.5.5.2 Case study 7: Electron paramagnetic resonance spectroscopy of Ti3+-Al and Th3+-Al bimetallics 1.07.5.6 Magnetism 1.07.5.6.1 Case study 8: Electronic structures of plutonium single molecule magnets 1.07.5.7 Mössbauer spectroscopy 1.07.5.7.1 Computational methods for Mössbauer spectroscopy 1.07.5.7.2 Case study 9: Mössbauer spectroscopy of (iPrPDI)Fe(N2)2 1.07.5.8 X-ray absorption spectroscopy (XAS) 1.07.5.8.1 Computational methods for X-ray absorption spectroscopy 1.07.5.8.2 Case study 10: Ni K-edge X-ray absorption spectroscopy of (iPr2NNF6)NiNO 1.07.5.8.3 Case study 11: Ligand K-edge X-ray absorption spectroscopy for evaluating lanthanide covalency 1.07.5.9 X-ray emission spectroscopy (XES) 1.07.5.9.1 Case study 12: Evaluating XES capabilities for probing NO coordination modes and reduction 1.07.6 Mechanism 1.07.6.1 Computational methods for calculating reaction mechanisms 1.07.6.1.1 Case study 13: Mechanism of CdCO2 bond formation at Cu, Rh and Pd 1.07.7 Current limitations and outlook 1.07.7.1 Case study 14: The ever elusive Grignard reaction Acknowledgments References

176

Comprehensive Organometallic Chemistry IV

177 177 177 178 178 178 179 179 180 180 180 180 181 182 182 182 183 183 185 185 185 185 186 186 186 187 187 188 189 189 190 190 190 191 191 192 193 193 193 194 194 195 195 196 197 199 199 200 200 201 201

https://doi.org/10.1016/B978-0-12-820206-7.00099-8

Computational Methods in Organometallic Chemistry

1.07.1

177

Introduction

Improvements in computer efficiency, cost, accessibility, and power have led to a vast increase in the accessibility of computations in a variety of chemical applications.1 Perhaps where computational chemistry has made the largest mark on organometallics is in facilitating determination of electronic structures. For organometallic chemistry, this enables computational methods to become part of the regular toolbox, even for chemists who have minimal formal training in computational chemistry. However, though the growing ease of computations makes them more routine for organometallic systems, the computational output can be interpreted incorrectly if proper approaches and techniques are not used. This is not to discourage synthetic chemists from learning computational methods, merely a word of caution.2 Hopefully this chapter enables those less familiar with computational methods to gain a greater understanding for communicating with collaborators, reading the literature, interpreting results, and beginning their own calculations. This chapter aims to give the reader a basic understanding of several computational methods commonly used in organometallic chemistry, showing the importance of benchmarking and offering practical examples through case studies that highlight specific applications and methods for problem solving. The chapter is geared toward experimentalists who are interested in learning about applications and current advances in computational chemistry for systems that they study. For the detailed computational methods, basis sets, and functionals, the reader should consult the papers cited within the chapter. For a more comprehensive discussion of computational methods and theory in organometallic chemistry, the reader is referred to textbooks on the topic.3–5

1.07.2

Density functional theory (DFT) and time-dependent density functional theory (TD-DFT)

Density functional theory (DFT) is in principle an exact theory that is widely applied in organometallic chemistry and is highly successful for calculating ground states of organometallic complexes.5–10 TD-DFT is the main method used to calculate excited states of organometallic systems.11 TD-DFT is based on the Runge-Gross Theorem, which assumes that there is a direct correlation between the electron densities at all points and a time-dependent external potential for a system with a given initial state.12 Even for calculating spectroscopy, time-dependent density functional theory (TD-DFT) often has relatively good agreement with experiment,10,13–17 although there are significant limitations.18 A general rule of thumb is that TD-DFT of higher energy spectroscopies (such as X-ray spectroscopies) has much better agreement with experiment than for lower energy spectroscopies (such as UV–vis). When calculated energies are large in magnitude, small uncertainties in energies have a minimal impact, but when energies are small in magnitude, small uncertainties in energy have a relatively larger effect on the accuracy of the result. Calculated intensities are proportional to the oscillator strength for a given transition between states and have no linewidth (often described as “sticks” in calculated spectra). Typically, line broadening is added to calculated spectra to offer an approximation of what would be observed experimentally. Current TD-DFT development includes creating new functionals to improve computational agreement with experiment, but most related work is in multireference and ab initio methods (see Section 1.07.3).17 One must ensure that a proper computational approach is taken when calculating excited states, however TD-DFT is likely to remain a go-to method for these types of calculations because it is relatively fast and cheap. Most DFT calculations are conducted using the Kohn-Sham formalism,19 so this will be the focus of this section.

1.07.2.1

Functional

The functional describes a relationship between energy and the electron density in a DFT calculation. The primary approximation in DFT is that the exchange and correlation effects are combined into one exchange-correlation (XC) energy, for which the exact solution is unknown, but is approximated by fits to experimental data or model systems to create XC functionals.1 The selection of the appropriate functional is critical to a successful computational outcome, where the three main types of functionals are local density approximations (LDA), generalized gradient approximations (GGA), and hybrid functionals. A full overview of the benefits and drawbacks of over 200 density functionals has been reported, 20,21 and an excellent discussion of functional choice may be found in “Which functional should I choose?”22 Briefly, LDAs are the most simple functionals, as they factor in only the density of a uniform electron gas at a point, but LDA calculations have errors on the order of 30 kcal mol−1 for molecules.23,24 GGAs factor in the density and the gradient at a point, and there are various functionals in this category depending on how the density and gradient are accounted for.25 GGAs require more computational power and time to calculate (often referred to as more “expensive” or “costly”), but they offer improvement in calculating transition state barriers and bond dissociation energies over LDA. A sub-class of GGAs are meta-GGAs which also factor in the kinetic energy density.26 Hybrid functionals start with a GGA or meta-GGA and add exact exchange from Hartree-Fock theory, resulting in the possibility of a much more accurate functional suited to a given system or property. Hybrid functionals are more expensive than GGAs, but generally feasible with current computational power for single point and geometry optimization calculations of small molecules (year 2021). BP86 is relatively commonly used because as a GGA functional it is faster than hybrid functionals, although it tends to overestimate metal-ligand covalency.27 This results in metal-ligand bond distances often optimizing to a distance that is shorter than experimentally observed by X-ray crystallography. The TPSSh functional is a hybrid functional (where TPSS is recognizable as a component in several meta-GGAs) that has also gained considerable use with fewer systematic errors and is less computationally

178

Computational Methods in Organometallic Chemistry

demanding than many other hybrid functionals, with less exchange than B3LYP.26,28 B3LYP is one of the most common hybrid functionals used in geometry optimizations for organometallic complexes because it offers a good balance between cost and accuracy.21 There is some discussion about the ideality of this functional,29 however given that so many studies have been conducted with it, it is useful to keep in mind when comparing systems. B3LYP is a variation on B3LYP with 15% Hartree-Fock (HF) exchange,30 as compared to 20% HF exchange in B3LYP, that can offer improved accuracy for calculating energetics. PBE0 is currently considered to be the most accurate for geometry optimizations of closed-shell systems across the periodic table.31 Some functionals have been developed specifically for or are well-tested for spectroscopic calculations with transition metals, so it is very important to use the same functional that was used in the initial calibration and benchmarking studies, unless an entirely new calibration is made. Additional computational parameters such as basis set, grid, corrections, and the such affect the global outcome of a calculation and must be tested in tandem with the functional in calibrations before production runs.

1.07.2.2

Basis set

A basis set refers to the functions (not to be confused with functional, Section 1.07.2.1) used to describe the wavefunction. Typically, the basis set consists of basis functions that approximate atomic orbitals, so the choice of basis set is highly dependent on the elements present in the calculation. The choice is also dependent on the type of calculation being conducted. Slater-type atomic orbitals (STOs) can be used for each electron in a system and are among the simplest basis sets.32 Gaussian-type orbitals (GTOs) use combinations of Gaussian functions to represent STOs and improve computational speed, even though a single Gaussian function does not have as accurate of a shape for an atomic orbital.33 GTOs are most commonly used for organometallic systems, as are large basis sets which offer higher flexibility with a large number of basis functions. Triple-z basis sets are combinations 3 of GTOs (or STOs) with different z that offer improved accuracy by allowing orbitals of the same type (such as px, py, pz) to differ in size to account for variations in bonding. Smaller basis sets resulting in lower accuracy but higher affordability include double-z basis sets that are combinations of two STOs or GTOs that have different z. Double-z and triple-z GTO basis sets are quite commonly used for organometallic chemistry. To facilitate the balance between expediency and cost, it is common to use a higher quality (such as triple-z) basis set for the metal center and any atoms immediately bound to the metal center, and a lower quality (such as double-z) basis set for all other atoms. Common combinations that are found in geometry optimizations from a range of papers include the triple-z def2-TZVP paired with the double-z def2-SV(P) Karlsruhe basis sets.34,35 Pople basis sets including the double-z 6-31G and triple-z 6-311G are also relatively common,36,37 but tend to be less consistent over a range of elements. Similarly, some basis sets have been optimized for specific spectroscopic calculations, so care must be taken to use the appropriate basis set for a given spectroscopic calculation. Most all-electron basis sets are only defined only for elements through Xe, so f-element calculations typically utilize either the LANL-type basis sets, or the Stuttgart-Dresden basis sets.38–40 Both of these utilize effective core potentials (ECP) to account for relativistic effects,39,41–43 although there is ongoing work to develop all-electron basis sets for actinides.44 An ECP is used to treat core electrons together, as opposed to explicitly calculating each one, which decreases the cost of the calculation and shortens computational time.

1.07.2.3

Additional considerations

Any time a calculation is conducted there are a range of other parameters that can be included. Once again, it is absolutely critical that the same parameters are used for a given set of calculations that one is interested in comparing. As soon as one component is changed, the results cannot be compared. Organometallic systems (especially spectroscopic calculations) often require careful selection of convergence settings and thresholds. For example, using the ORCA program it is common to include a slow convergence to ensure an appropriate minimum is obtained, along with tight convergence criteria, convergence energy change of 10−7 Eh (Hartree), and residual error thresholds of 10−6 Eh. If convergence criteria are too loose or large, a proper minimum may not be found, but if they are too tight or small, a calculation may never converge. RIJCOSX is an approximation that can speed up calculations of organometallics that use hybrid functionals and RI is used for GGA and meta-GGA functionals.45,46 Benchmarking calculated energies with experimental data is discussed in Section 1.07.5.1. Solvents can be accounted for by using a polarization continuum model (PCM) such as COSMO.47 Dispersion corrections are also often necessary to account for a van der Waals-type attraction between atoms and molecules that are not directly bound to each other (often dispersion and van der Waals are used interchangeably).48 These are especially important in large systems.

1.07.2.4

Beginning calculations

The three ideal pieces of information an experimentalist should have before embarking on computational work is to know the overall spin state of the molecule, the overall charge of the molecule, and to have a reasonable 3-D structural input (ideally from a crystal structure). These components are included in the computational input file, which is the file that tells the chosen computational program what to do (referred to as “inputs” or “computational inputs” in this chapter). The overall spin state can be determined from the effective magnetic moment using a magnetic susceptibility balance (MSB), superconducting quantum interference device (SQUID) magnetometer, or by NMR using the Evans method,49,50 for example. Structural inputs are generally in the form of xyz coordinates that define the positions of each atom. Notice that atom positions do not infer anything about the

Computational Methods in Organometallic Chemistry

179

presence of bonds; they simply describe where the nucleus of each atom is. If the xyz coordinates are taken from a crystallographically determined structure, it is critical to check if there is additional solvent in the structure. Typically, the solvent molecule(s) is removed for calculations. Without these experimental parameters, the potential number of computational inputs are vast and would be very time consuming to conduct, with high potential for error. For example, if the spin state of a proposed Fe(II) molecule is unknown, the arrangement of 6 d-electrons in 5 d-orbitals could result in overall spin states of S ¼ 2, S ¼ 1 and S ¼ 0. This would require at least 3 calculations to be conducted, and most likely multiple inputs would be needed for each spin state for geometry optimizations and spectroscopic properties. This would result in up to 6–9 calculations to be conducted and analyzed, and if the ligand has any potential for being redox-active, even more calculations would need to be conducted and analyzed (see Section 1.07.2.7). By contrast, if the spin state is known, only 1–3 calculations would need to be conducted. This example demonstrates how planning and some experimental data can greatly improve the efficiency of computational work. If a structure of a molecule is unknown, it is best to begin with a structure that was crystallographically determined for a structurally similar molecule. Several programs are available for modifying structures (see Section 1.07.2.4.3), but it is difficult to build new structures in 3-D space that offer a chemically reasonable starting point for most computational programs. If the structure deviates too far from what is chemically reasonable (such as very long interatomic distances), it is likely that no solution (or a non-realistic solution) to the self-consistent field (SCF) will be found. Modifications that are relatively trivial include removing solvent and changing atom types, such as modifying the 3-D structure of a palladium phosphine complex and replacing palladium with nickel. Replacing a fragment of a ligand with a lower nuclearity fragment (such as tert-butyl to methyl) is a process known as truncation that is relatively straightforward and can speed up calculations because fewer basis functions need to be included. More complicated modifications can include taking a ligand from one structure and combining it with another structure to build a new geometry to use as the input for new calculations.

1.07.2.4.1

Geometries and geometry optimization

In a geometry optimization, the positions of the atoms are varied in small increments to determine the lowest energy structure, through an iterative process that minimizes the energy gradient (force) on each atom.1 Generally, it is prudent to conduct a geometry optimization for a molecule before embarking on more sophisticated electronic structure or spectroscopic calculations. Note that crystallographic structures have estimated standard deviations in bond lengths of 0.001 to 0.01 A˚ , which are often an underestimation of the actual uncertainty.51 Since an estimated standard deviations represents only 68% certainty limit and there are dozens or hundreds of bond lengths in a molecule, it is likely that many of the experimentally determined distances have significant inaccuracies that could have important effects on the computational results. Another important consideration is that X-ray crystallography does not determine the positions of nuclei (the xyz positions of atoms in a calculation), but instead gives the positions of maximum electron density. For most atoms, the predominance of core electrons makes the centroid of electron density so close to the nucleus that the difference can be ignored. However, for hydrogen atoms, the only electron is engaged in bonding, and as a result X-ray crystal structures systematically underestimate the length of bonds to hydrogen by about 0.1 A˚ . This systematic error can cause substantial inaccuracies in computational results if geometry optimization is not performed. Finally, it is common (but becoming less so) for supporting ligands to be truncated before embarking on computations, in order to decrease the time for geometry optimization. However, there are many cases in which this truncation can alter the results. Given the increasing computational power available to the chemist, it is recommended to avoid truncation, and if it is used then it is important to perform benchmarking tests to assess any impacts of the truncation. Although there is some debate in the community, the most common functionals used for geometry optimization of organometallic systems are BP86 and B3LYP52–54 or PBE0,31 along with basis sets of triple-z quality for the metal center and the directly coordinating atoms, such as def2-TZVP with auxiliary basis sets.34,55–57 With these combinations, errors in relative energies are generally on the order 3 kcal mol−1 or less. It is important to perform a frequency calculation to ensure that the geometry is actually at a minimum, which is displayed by having only positive frequencies (or negative ones less than | n | < 10 cm−1). A few negative frequencies may be the result of an integration grid that is too small, or using an RIJCOSX approximation.58 Negative frequencies indicate that the calculation is at a saddle point and not at a minimum of the potential energy surface, which can be fixed by changing the input geometry a bit and restarting the calculation.58 The optimized geometry usually does not have a large dependence on functional, and so it is common to optimize a geometry using a simpler functional or basis set, and to “fix” this geometry for more expensive methods in single point calculations to calculate electronic or spectroscopic properties. Perhaps the easiest initial check when examining the results of a geometry optimization is to first view the resulting geometry (xyz file for example) in a molecular viewing program (such as Avogadro, ChemCraft or GaussView). This allows one to quickly visualize the structure and assess if there are any glaring errors. This is especially important to do if a calculation does not converge. It is very easy to accidentally omit H-atoms, or to have forgotten to remove solvent from a crystal structure, for example.

1.07.2.4.2

Single point calculation

In a single point calculation, the positions of the atoms are not varied and the total energy of the system is calculated.1 For each spectroscopy calculation, for example, a single point calculation is conducted. In cases where a crystal structure exists for a molecule and time is limited, a single point calculation without prior geometry optimization is sometimes used. Though this contradicts the advice in the previous section, the use of crystallographic coordinates typically results in reasonable agreement with experiment

180

Computational Methods in Organometallic Chemistry

when the desired output involves only the properties of the metal center (for example Mössbauer or XAS calculations). However, if possible it is desirable to first optimize the geometry to ensure the best quantum chemical description, especially when energies and vibrational frequencies are under investigation.

1.07.2.4.3

Software and common programs

Open source software or programs that are free for academic use include the ORCA quantum chemistry program package for conducting DFT, multireference, many-body perturbation, coupled cluster, and semi-empirical calculations,59 Quantum Espresso,60 and GAMESS.61 ORCA has a useful website and an active user forum that gives a lot of practical advice for conducting calculations.58 For visualization of orbitals and structures, Chimera offers a nice user interface,62 and Avogadro is a program that can be used to modify input structures and has some ORCA capabilities.63 VMD can also be used for visualizing orbitals and structures, with the additional capability of visualizing trajectories (such as IR frequencies).64–71 Some common commercial programs that are used in computational chemistry for organometallic systems include Gaussian,72 Molpro,73 and ADF74 for quantum chemical calculations, and ChemCraft for building molecules, visualizations, and viewing trajectories.

1.07.2.5

Closed Shell systems and restricted Kohn-Sham (RKS)

Closed shell systems, which are common in complexes of 2nd and 3rd row transition metals, are some of the organometallic systems that DFT can calculate most easily and accurately.11 Closed shell systems have a total spin of S ¼ 0, such that electrons can be treated as electron pairs and orbitals are treated as either empty (no electrons), or filled (two electrons). The computational approach for this is called the restricted Kohn-Sham (RKS) method.19

1.07.2.6

Open Shell systems and unrestricted Kohn-Sham (UKS)

As soon as there is one unpaired electron in the system, as is common for many first row transition metal complexes, even paired valence electrons must be treated individually because their interactions with the unpaired electron are different.11 The computational approach for this is called the unrestricted Kohn-Sham (UKS) method, and can be utilized for both closed shell (S ¼ 0) and open shell (S > 0) systems. Chemically, the computational input takes traditional two-electron orbitals with interacting electrons that are now split into one-electron orbitals where each electron has its own orbital. Spin up electrons are treated separately from spin down electrons in different spatial orbitals with differing energies. These one-electron orbitals are either empty (no electrons), or filled (1 electron), and are classified as alpha (the higher-population spin) or beta (lower-population spin). The pictures of one-electron orbitals deviate from traditional molecular orbital diagrams with two-electron orbitals, but are usually brought into this model using a method to find the pairing system that gives the maximum overlap between pairs of electrons (each pair being one alpha and one beta). The resulting pictures of quasirestricted orbitals are thus amenable to a traditional interpretation, but it is important to consult the output that gives the overlap (a percentage confusingly called S, which is a different S than the spin state. The term Soverlap will be used here.) between the corresponding pair of an alpha and a beta orbital. There is some ambiguity about the minimum Soverlap for a corresponding pair that indicates that the two electrons are truly paired, but certainly overlaps >95% can be treated as electron pairs. For a true closed-shell system, the UKS and RKS calculation solutions are effectively the same. Corresponding pairs of filled alpha and beta electrons with overlaps 2.214 The ZFS is a tensor that can be reduced to axial (D) and rhombic (E) ZFS parameters. The exchange coupling constant (Jex or isotropic interaction parameter) refers to a difference in energy between two magnetic alignments. For example, two coupled S ¼ ½ Cu(II) centers give two states, one with S ¼ 0 and another with S ¼ 1. If the S ¼ 0 state is lower in energy then it is described as antiferromagnetic coupling, and if the S ¼ 1 state is lower in energy then it is described as ferromagnetic coupling. Jex is related to the difference in energy between the two states (generally expressed as a positive number for ferromagnetic coupling and a negative number for antiferromagnetic coupling). Jex can be calculated using broken symmetry (Section 1.07.2.7), single determinant (SD),215 or a spin decontamination method.216 For systems with large Jex (>|100| cm−1) DFT is relatively good,8 but the main problem arises when there are multiple configurations that have small splitting and are nearly degenerate. In iron compounds, Jex values can easily range from +10 to −100 cm−1.217 It can also be difficult to distinguish D and Jex experimentally, since each of these can give small field-independent splitting of energy levels. As a result, accurate Jex measurements are usually only possible for systems with no ZFS.214 For hyperfine interactions, DFT typically makes predictions of dipolar couplings and contributions to the hyperfine coupling from the g-tensor and SOC that are too small.217 Single molecule magnets (SMMs) are an emerging area for quantum computing research,218–222 however computational design and prediction for SMM behavior is still under development. Currently, it is not well understood how to predict magnetization relaxation,223–225 and multiconfigurational approaches are necessary to accurately calculate magnetic properties in systems with multiple nearly degenerate configurations, and the electronic structures of many SMMs.

1.07.5.6.1

Case study 8: Electronic structures of plutonium single molecule magnets

One of the first reports of quantum chemical studies of transuranic SMMs compared the electronic structures of Tp3Pu [Tp− ¼ hydrotris(pyrazolyl)borate) and the analogous structure with a carbene ligand (Fig. 10).226 Actinides are of particular interest for SMMs because of their large spin-orbit coupling, however computational challenges for actinides make prediction and design of SMMs even more complicated. The geometry was optimized with BP86,165 TZ2P basis set for Pu and DZP basis set for other atoms. The ZORA approximation was added to account for relativistic effects.203,204 For Tp3Pu, both CASSCF and CASPT2 approaches with spin-orbit coupling were probed for the Pu(III) system that has 5 unpaired electrons in f-orbitals with a ground state of S ¼ 5/2 and 6H5/2.227 A CASSCF(5,7) calculation was conducted that included all roots (possible multiplicities) of 21 sextets, 224 quartets, and 490 doublets. The sheer number of 735 possible states would be extremely time intensive to analyze. However, plotting the energies of the calculated states based on multiplicity (Fig. 11) shows that the quartet and doublet states are so high in energy as compared to experimental 5f to 6d and LMCT transitions at 58283 cm−1 and 105,235 cm−1, respectively, that doublet and quartet 5f to 6d and LMCT transitions are unlikely to undergo spin-orbit coupling with the sextet states. Therefore, only transitions from sextet states were used in further calculations, bringing the total number of calculated states down to 21. This is a great example of where a quick evaluation and assessment of computational results with experimental data can allow for more efficient analysis. RASSI calculations accounted for spin-orbit coupling, resulting in a calculated first excited Kramers’ doublet of 373 cm−1, which is in reasonable agreement with the experimentally reported value of 332 cm−1 for the first excited state. A larger magnitude for this first excited state may correlate to a higher relaxation barrier. The magnetic susceptibility curve was also calculated to be in reasonable agreement with experiment, and calculations of the carbene-substituted complex suggested that it may have an even higher barrier for relaxation than Tp3Pu.

192

Computational Methods in Organometallic Chemistry

Fig. 10 Drawing (top) and 3-D representation (bottom) of Tp3Pu (left) and the analogous structure with a carbene ligand. Reprinted with permission from Gaggioli, C. A.; Gagliardi, L. Theoretical Investigation of Plutonium-Based Single-Molecule Magnets. Inorg. Chem. 2018, 57(14), 8098–8105, https://doi.org/10. 1021/acs.inorgchem.8b00170. Copyright 2018 American Chemical Society.

Fig. 11 Relative energy ranges calculated using CASSCF of sextet, quartet and double states for Tp3Pu. The green and black lines correspond to the first Pu 5f to 6d transition at 58283 cm−1 and ligand to metal charge transfer at 105235 cm−1, respectively. Reprinted with permission from Gaggioli, C. A.; Gagliardi, L. Theoretical Investigation of Plutonium-Based Single-Molecule Magnets. Inorg. Chem. 2018, 57(14), 8098–8105, https://doi.org/10.1021/acs. inorgchem.8b00170. Copyright 2018 American Chemical Society.

1.07.5.7

Mössbauer spectroscopy

The Mössbauer effect is the recoilless nuclear resonance absorption of a gamma ray by a nucleus.228–230 A good Mössbauer isotope needs to have a low excited state, long life-time (line width), reasonable source and reasonable hyperfine splitting (nuclear spin). 57 Fe is the most common nucleus used for Mössbauer spectroscopy, and spectra are sensitive to iron oxidation state, iron spin state, coordination environment, covalency, molecular symmetry, and magnetic interactions.217,231 The most common experiment involves a 57Co source that generates gamma rays of the correct energy to excite 57Fe, which has a natural abundance of 2.2%. For solid organometallic complexes, this natural abundance is sufficient to collect a good Mössbauer spectrum, while in solution samples or metalloproteins it is often necessary to use 57Fe enriched samples.

Computational Methods in Organometallic Chemistry

1.07.5.7.1

193

Computational methods for Mössbauer spectroscopy

Computational agreement between experimental and DFT calculated isomer shifts (d) and quadrupole splittings (DEQ) for Fe systems is usually good, with mean average errors on the order of 0.10 mm s−1 for the isomer shift (d) and 0.50 mm s−1 for the quadrupole splitting (DEQ).232–234 A more recent calibration reported accuracies of 0.13 mm s−1 for the isomer shift (d) and an improved accuracy of 0.36 mm s−1 for the quadrupole splitting (DEQ).235 Practically speaking, these error ranges usually mean that the isomer shift can be used as a benchmark for testing a computational electronic structure model, while one needs to be more careful with benchmarking to the quadrupole splitting. Calculated parameters are suitable for comparison to experimental data up to liquid nitrogen temperatures of 77 K where contributions from thermal motion are around 0.02 mm s−1.235,236 By contrast, thermal motion contributes −0.1 mm s−1 at room temperature. It should also be noted that most in-house Mössbauer instruments used by organometallic chemists are able to experimentally determine only the absolute value of the quadrupole splitting. Application of large fields (>1 T) causes splitting, and these “magnetic Mössbauer” experiments give the sign of the quadrupole splitting, which offers a more discerning experimental comparison for benchmarking a computational model.

1.07.5.7.2

Case study 9: Mössbauer spectroscopy of (iPrPDI)Fe(N2)2

An example of the need for multiple computational inputs and benchmarking to Mössbauer spectroscopy was reported for the bis(imino)pyridine iron dinitrogen complex, (iPrPDI)Fe(N2)2 (Fig. 12).79 This complex is diamagnetic, however the ligand bond distances suggest an electronic structure between (iPrPDI)1− and (iPrPDI)2−. Therefore, computational inputs of BS(1,1) and BS(2,2) were reasonable to test, which would describe (iPrPDI)1− FeI(N2)2 (low-spin iron(I)) and (iPrPDI)2− FeII(N2)2 (intermediate-spin iron(II)), respectively. For comparison, the RKS input would correspond to (iPrPDI)Fe0(N2)2 with low-spin iron(0). After geometry optimizations, the BS(1,1) and BS(2,2) solutions converged to the same BS(1,1) solution that was 2.4 kcal mol1− lower in energy than the distinct RKS solution. Based on energetics alone, these two possibilities could not be distinguished, since they are within the error of approximately 3 kcal mol1−. Both solutions had good agreement with calculated bond distances and bond angles. For this reason, it was critical to have an additional experimental test to distinguish calculated electronic structure possibilities. A magnetic Mössbauer spectrum yielded d ¼ 0.39 mm s−1 and DEQ ¼ −0.53 mm s−1, which could be compared to calculated values of d ¼ 0.38 mm s−1 and DEQ ¼ −0.65 mm s−1 for the RKS solution, and d ¼ 0.51 mm s−1 and DEQ ¼ −0.81 mm s−1 for the BS(1,1) solution. Based on the isomer shifts, the RKS solution of d ¼ 0.38 mm s−1 was well within the error of 0.10 mm s−1 as compared to the experimental value of d ¼ 0.39 mm s−1, while the BS(1,1) solution of d ¼ 0.51 mm s−1 was outside the error. Based on isomer shift alone, the RKS solution was determined to be the better electronic structure description of the system, supporting a description as an Fe(0) complex. This is an excellent example of using experimental data to determine the validity of a calculated electronic structure. Relative calculated energies alone can offer some support for the structure with the calculated lower energy, however an experimental comparison, such as good agreement between experimental and calculated Mössbauer parameters is much stronger support for a given computational model.

1.07.5.8

X-ray absorption spectroscopy (XAS)

X-ray absorption spectroscopy (XAS) involves excitation of a core electron into empty valence orbitals with possible transitions from electrons in the first shell (K-edge, n ¼ 1), second shell (L-edge, n ¼ 2), and third shell (M-edge, n ¼ 3).237 For organometallic complexes, the K-edge is typically the most useful because it involves a 1 s to nd transition that is often highly sensitive to metal oxidation state. For this reason, the K-edge is often used to determine the oxidation state of first row transition metals, in particular, as shifts of approximately 1 eV (1240 nm) of the 1 s to 3d transition correspond to 1-electron changes of oxidation state at the metal center.237–239 Metal K-edge spectroscopy is also used for collecting extended X-ray absorption fine structure (EXAFS) data that yield 3-D structural information about systems for which crystal structures cannot be obtained.237,239 The most common applications for EXAFS in organometallic chemistry are for bioinorganic systems, and unstable intermediates. Other applications of XAS that gained interest in recent years for the organometallic community are ligand K-edges for evaluating ligand oxidation states and covalency, 240–246 as well as significant advances in XAS of f-elements.145,147,243,247,248 A challenge with the f-elements is that the metal K-edge has very high energy, so lower energy M- or L-edges are more experimentally feasible but can be more challenging to understand

Fig. 12 Structure of (iPrPDI)Fe(N2)2.

194

Computational Methods in Organometallic Chemistry

(Section 1.07.4.2.1). Ligand K-edge or L-edge XAS overcomes some of these challenges by probing interactions between the metal and the ligand from the perspective of the ligand instead, giving complementary insights into bonding and electronic structure (see Section 1.07.5.8.3).

1.07.5.8.1

Computational methods for X-ray absorption spectroscopy

The currently accepted method for calculating XAS is using time-dependent density functional theory (TDDFT).249–251 Computational resources are now sufficiently fast and inexpensive that the more time-intensive TDDFT calculations do not cause undue burden, however it is worth noting that simple ground state model calculations of ligand field multiplet models of XAS (Section 1.07.4) are oftentimes in surprisingly good agreement with experiment.238,252,253 TD-DFT calculates transitions as “sticks” without any linewidth, so the calculated spectra are broadened to create an envelope that is more easily compared to experimental data. This broadening should be selected to be on the order of what is observed experimentally (typically 1.5 eV for a first row transition metal K-edge). Each stick corresponds to a transition, and in the calculation the orbitals contributing to these transitions can be evaluated. Calibrations are critical for interpretation of XAS spectra, and the calibration is different for each particular element. The calculated spectra typically have to be shifted to match experimental energies, and it is most rigorous to shift all calculated spectra by the same constant for a given method and element to ensure no bias in the calculation. Changing the functional used to calculate a spectrum results in a change in the energy shift, requiring a new calibration. For this reason, it can sometimes be beneficial to select a basis set and functional to fit a known calibration. Extensive XAS calibrations have been reported for Fe,249,254 for example with the BP86 functional52,255 with CP(PPP) basis set.256 Recent advances include using a combined ROCIS/DFT approach, which allows metal K-edge, L-edge, and M-edge XAS spectra to be calculated for systems of over 700 atoms and clusters with over 50 metal centers.257

1.07.5.8.2

Case study 10: Ni K-edge X-ray absorption spectroscopy of (iPr2NNF6)NiNO

An example using X-ray absorption (XAS) spectroscopy for testing electronic structures is in the reduction of end-on (iPr2NNF6) NiNO to [(iPr2NNF6)NiNO][K].87 In Enemark-Feltham notation,130 (iPr2NNF6)NiNO is {NidNO}10 while [(iPr2NNF6)NiNO][K] is {NidNO}11. Since the nacnac ligand (iPr2NNF6) is an anionic ligand, the Ni center in (iPr2NNF6)NiNO is either Ni(I) with a neutral NO, or Ni(II) with anionic NO. Especially in these systems, it is then critical to have additional experimental benchmarks for evaluating a given computational model and determining an electronic structure. The experimental XAS (Fig. 13) have edge energies consistent with Ni(II) metal centers, with the reduced complex appearing slightly oxidized at the nickel center by virtue of a slight shift of the pre-edge (8332 eV, 0.15 nm) to higher energy. Geometry optimizations were conducted with B3LYP and def2-TZVP basis sets on the Ni, N, and O, with def2-SVP on all other atoms. The lowest energy solution was a BS(2,2) solution corresponding to (iPr2NNF6)NiII(NO)1− with an intermediate spin S ¼ 1 Ni(II) antiferromagnetically coupled to S ¼ 1 NO1−. A UKS S ¼ ½ solution

Fig. 13 Experimental (left) and TD-DFT calculated Ni K-edge XAS spectra for linear/end-on (iPr2NNF6)NiNO, linear/end-on [(iPr2NNF6)NiNO][K], and side-on [(iPr2NNF6)NiNO][K]. Reprinted with permission from Kundu, S.; Phu, P. N.; Ghosh, P.; Kozimor, S. A.; Bertke, J. A.; Stieber, S. C. E.; Warren, T. H. Nitrosyl Linkage Isomers: NO Coupling to N 2 O at a Mononuclear Site. J. Am. Chem. Soc. 2019, 141(4), 1415–1419, https://doi.org/10.1021/jacs.8b09769. Copyright 2019 American Chemical Society.

Computational Methods in Organometallic Chemistry

195

for [(iPr2NNF6)NiII(NO)2−][K] was lowest in energy, corresponding to low spin S ¼ 0 Ni(II) with the spin density on NO2−. These solutions were benchmarked to EPR data and crystallographic data. To additionally test the validity of this electronic structure, TD-DFT XAS calculations were compared to experimental XAS data. TD-DFT calculations were conducted with the BP86 functional52,255 with CP(PPP) basis set256 on the metal and TZVP56 on other atoms with a COSMO solvation model,47 as described in a previous Fe XAS calibration paper.254 Broadening of 1.5 eV resulted in the calculated spectrum in Fig. 13. The TDDFT spectrum had good agreement to the experimental XAS spectrum with respect to relative energies and intensities. The feature at 8335 eV has higher experimental and calculated intensity for [(iPr2NNF6)Ni(NO)][K] than for (iPr2NNF6)NiNO, and is assigned to be a Ni to ligand p transition based on the orbital contributions to the sticks. These results supported overall reduction of NO since a decrease in intensity of the observed feature is consistent with the p orbital having an additional electron (ie. being filled). In combination with DFT, the TD-DFT calculations offered comparison to experiment to further support the electronic structure description. Although ab initio methods could potentially offer additional insight toward the bonding of the system, the reported combination was sufficient for assessing metal-based versus NO-based reduction.

1.07.5.8.3

Case study 11: Ligand K-edge X-ray absorption spectroscopy for evaluating lanthanide covalency

When X-ray absorption spectroscopy is used to study lanthanide and actinide systems, the L3,2-edge (2p to 3d transitions) or M5,4-edge (3d to 4f transitions) are usually probed, because these bands are in an experimentally achievable energy range. However, these spectra are often dominated by multiplet interactions resulting from both initial and final states having multiple possible electron configurations. Numerous recent studies have also demonstrated the feasibility of using ligand K-edge XAS for evaluating covalency of lanthanides and f-elements in combination with computational studies.145,147,247,248,258,259 This has offered a new experimental avenue for benchmarking f-element calculations and highlights that a simple ionic description of lanthanides is inaccurate. In the current case study, Cl K-edge XAS studies were conducted for lanthanide hexahalide complexes, LnCl3− 6 (Ln ¼ Ce, Nd, Cm, 145 Eu, Gd) and CeCl2− The intensity mechanism for ligand K-edge XAS results from electric dipole allowed transitions,250 such that 6 . observed intensities must result from mixing of Cl 3p orbitals with Ln 4f or 5d orbitals. Higher intensity features result from an increase in Cl 3p character in bonding. For these octahedral complexes, the 4f orbitals with t1u and t2u symmetry and the 5d orbitals of t2g and eg symmetry have the correct symmetry to be observed in the Cl K-edge XAS spectrum. TD-DFT calculations for the Cl K-edge were used to assign transitions from the Cl 1s to 3p orbitals. TD-DFT calculations from Cl 1s orbitals to the valence shell were in excellent agreement with experiment (Fig. 14), allowing for contributions of 4f versus 5d bonding to be elucidated. In Fig. 14, the yellow lines are the calculated transitions or “sticks,” for which the orbitals involved in the transition can be examined. This allows each “stick” to be evaluated and assigned to valence orbitals based on which orbitals contribute to the transition. Since the Cl 3p orbitals mix with metal orbitals, these can be evaluated based on metal contributions such as the 4f or 5d orbitals. The colored envelopes (pink, blue, and green) are broadened calculated spectra based on the “sticks” and are colored based on the metal orbital assignment. Especially for CeIII and CeIV, the broadened calculated spectra agree very closely with experimental data (black), both with respect to energy and intensity. Results show that lower energy features result from transitions to 4f orbitals, followed by transitions to 5d p-orbitals and finally 5d s-orbitals. Unfortunately, the resolution of the Cl K-edge is not sufficient to observe 4f-4f orbital splitting, which was predicted to increase in the order CeIII < NdIII < SmIII < EuIII because of multiplet effects. Fitting of the experimental spectra supported 7–13% of Cl 3p character per bond with 5d t2g orbitals, with an overall decrease in mixing from CeIII > NdIII > SmIII > EuIII > GdIII. Remember, that Cl 3p mixing is the intensity mechanism for Cl K-edge XAS. Neither DFT nor TD-DFT calculations found significant Cl 3p character per bond with 5d t2g orbitals (0–1%), suggesting a possible limitation of these computational methods and the need for more sophisticated SOC and multiconfigurational calculations.

1.07.5.9

X-ray emission spectroscopy (XES)

A spectroscopic method that is still in the development stages for characterizing organometallic systems is X-ray emission spectroscopy (XES).260–262 XES involves detection of photons emitted during relaxation of valence electrons following excitation of a 1s electron, which allows the energies of filled orbitals to be probed. For organometallic chemistry, the most common XES experiment is metal Kb XES, where 3p to 1s (Kb1,3) and ligand-centered ns/np to metal 1s (Kb” and Kb2,5) transitions are probed. In particular, the Kb” and Kb2,5 or valence-to-core (VtC) regions in spectra are dominated by transitions from orbitals having primarily ligand character, such that the VtC offers unique insights toward bonding between the metal and ligands. Already, the technique has been applied in several metalloprotein systems,263–266 and has also been shown to be sensitive to metal spin state,267,268 ligand hybridization,269 ligand protonation,270,271 ligand bond angles and geometry,272–274 ligand activation,272,275–277 number of bound ligands,79,269 and ligand identity.278–283 XES is perhaps most well-known to the bioinorganic community for the role in determining the central atom in the FeMoCo cluster of nitrogenase to be carbon, as opposed to nitrogen or oxygen.278 For experimentalists, this highlighted an important advantage over XAS for determining structures of intermediates that cannot be crystallographically characterized, since XAS cannot distinguish similar light atoms such as C, N, and O.237 Computationally, the firmness of this conclusion was dependent on the benchmarking in previous papers describing XES on model complexes that were crystallographically characterized in combination with DFT calculations to determine spectral contributions and experimental limits of XES. Model systems included ferric and ferrous coordination complexes,269 iron clusters,279 and iron dinitrogen complexes.79,284

196

Computational Methods in Organometallic Chemistry

2 Fig. 14 Experimental (black) versus TDDFT calculated (pink ¼ 4f, blue ¼ 5d, green ¼ 5d) Cl K-edge XAS spectra for LnCl3− 6 (Ln ¼ Ce, Nd, Sm, Eu, Gd) and CeCl6 . Reprinted with permission from Löble, M. W.; Keith, J. M.; Altman, A. B.; Stieber, S. C. E.; Batista, E. R.; Boland, K. S.; Conradson, S. D.; Clark, D. L.; Lezama Pacheco, J.; Kozimor, S. A.; Martin, R. L.; Minasian, S. G.; Olson, A. C.; Scott, B. L.; Shuh, D. K.; Tyliszczak, T.; Wilkerson, M. P.; Zehnder, R. A. Covalency in Lanthanides. An X-Ray Absorption Spectroscopy and Density Functional Theory Study of LnCl6 x – ( x ¼ 3, 2). J. Am. Chem. Soc. 2015, 137 (7), 2506–2523, doi: 10.1021/ja510067v. Copyright 2015 American Chemical Society. −

The computational method that is still most widely used is a ground state DFT calculation and one-electron approximation that effectively takes into account differences in orbital energies to calculate possible transitions.269 While it is not the most physically accurate description, the computational agreement with experimental data for model complexes has been excellent for a range of systems based on Cr,281,283 Mn,282,285–289 Fe,79,275,279,280,284 Co,268 and Cu.265 The BP86 functional52,255 with CP(PPP) basis set256 on the metal and TZVP56 on other atoms with a COSMO solvation model47 is quite common.269 Calculated transitions or sticks for first row transition metal Kb XES are generally broadened by 2.5 eV (496 nm) to match experimental resolution limits.269 Current areas of development include using model complexes with other metal centers to establish experimental correlations to computations, theory development for resonant X-ray emission (RXES or RIXS),290 theory development for the Kb1,3 region, as well as development of time-resolved XES methods for characterizing transient organometallic intermediates in reactions. For organometallic chemists, this offers potentially a new probe for in situ characterization of reactions.

1.07.5.9.1

Case study 12: Evaluating XES capabilities for probing NO coordination modes and reduction

An example of how DFT can be used to understand spectral contributions to XES is highlighted in the same (iPr2NNF6)NiNO and [(iPr2NNF6)NiNO][K] complexes from Section 1.07.5.8.2.272 These complexes were selected since they are fully crystallographically characterized, spectroscopically characterized with EPR, IR, and XAS data, and electronic structures were established by DFT.87 The three complexes used for XES studies were end-on (iPr2NNF6)NiII(NO)1−, end-on [(iPr2NNF6)NiII(NO)2−][K] and side-on [(iPr2NNF6)NiII(NO)2−][K]. Sensitivity of XES to NO reduction was probed since the complexes span (NO)1− and (NO)2− oxidation states, while the Ni oxidation state remained the same Ni(II). NO coordination modes could be tested since [(iPr2NNF6)NiII(NO)2−] [K] complexes were isolated with both side-on and end-on NO coordination. When using DFT to learn more about XES spectral features, it is critical to use experimental systems that are very well characterized as model systems before embarking on unknowns.

Computational Methods in Organometallic Chemistry

197

Fig. 15 Experimental (left) and calculated (right) Ni Kb VtC XES for end-on Ni-(NO1−), end-on Ni-(NO2−), and side-on Ni-(NO2−). A shift of 217.0 eV and broadening of 2.5 eV were applied to calculated spectra. Reprinted with permission from Phu, P. N.; Gutierrez, C. E.; Kundu, S.; Sokaras, D.; Kroll, T.; Warren, T. H.; Stieber, S. C. E. Quantification of Ni–N–O Bond Angles and NO Activation by X-Ray Emission Spectroscopy. Inorg. Chem. 2021, 60(2), 736–744. Copyright 2021 American Chemical Society.

Experimental Kb XES for these complexes were compared to computational XES using the BP86 functional52,255 with CP(PPP) basis set256 on the metal and TZVP56 on other atoms with a COSMO solvation model.47 The experimental and calculated XES spectra were similar, both with respect to calculated intensities and energies (Fig. 15). Notably, the XES spectra for all three complexes were distinct, suggesting that XES can be used to probe 1-electron reduction events at NO, and different NO coordination modes. Further examining the orbital contributions to spectral features showed that signal A in Fig. 15 corresponds to the NO s bonding orbital, which gains intensity for side-on NO coordination. Signal C in Fig. 15 corresponds to the NO s (antibonding) orbital, which decreases in intensity upon NO reduction and decreases in intensity for side-on NO coordination. The utility of DFT for understanding spectral features is in the ability to calculate structures that cannot be experimentally probed. Once the experimental and calculated XES spectra were compared and found to be in good agreement, the structures of the complexes were altered to probe the effect of lengthening the NdO bond, lengthening the NidN bond, and changing the NidNdO angle. The calculated spectrum for changing the NidNdO angle from 80 to 160 is presented in Fig. 16. The calculations indicate that the feature corresponding to the NO s bonding orbital gains intensity for decreasing NidNdO angles (toward side-on coordination). By contrast, the signal corresponding to the NO s (antibonding) orbital loses intensity for decreasing NidNdO angles (toward side-on coordination). This suggests that XES may be a useful experimental tool for probing NO coordination modes. Finally, this method for calculating XES was used to probe the possible utility for using XES to characterize the copper nitrite reductase (NiR) active site (Fig. 17). Input structures for the active site were taken from reported crystal structures,291 which are rare examples of crystallographic characterization of possible NiR intermediates. Cu Kb VtC XES calculations of these structures display similar trends to the calculations of the NidNO complexes. The signal at lowest energy, has primary contributions from the 2s/NO s bonding orbital, which has the highest intensity for side-on NO coordination. By contrast, the feature corresponding to the 2s/NO s orbital has highest intensity for end-on NO coordination. This is exactly the same trend observed in the NidNO complexes, and suggests that these trends may be applicable over a range of metals and ligand systems. Overall, this study highlights how DFT can be used to understand spectroscopic methods and generate new possibilities for characterizing complex systems.

1.07.6

Mechanism

Organometallic mechanisms present more of a challenge for computational chemistry because of the additional variables such as intermolecular interactions, solvents, and the need to query parts of the potential energy surface far from the equilibrium geometry. For a more detailed discussion, the reader is pointed to an excellent paper highlighting common challenges and components to be aware of.292 Two of the most common pitfalls for mechanistic studies result from incorrect use of simplified structures and errors in technical aspects of conducting quantum chemical calculations.292 Many times systems are truncated (such as through replacement of ligands with lower nuclearity analogues) to increase the efficiency of the calculation, but care must be taken to not oversimplify

198

Computational Methods in Organometallic Chemistry

Fig. 16 Calculated Ni Kb VtC XES for Ni-(NO2−) with NidNdO angles ranging from 80 to 160 . A shift of 217.0 eV and broadening of 2.5 eV were applied to calculated spectra. Reprinted with permission from Phu, P. N.; Gutierrez, C. E.; Kundu, S.; Sokaras, D.; Kroll, T.; Warren, T. H.; Stieber, S. C. E. Quantification of Ni–N–O Bond Angles and NO Activation by X-Ray Emission Spectroscopy. Inorg. Chem. 2021, 60(2), 736–744. Copyright 2021 American Chemical Society.

Fig. 17 Calculated Cu Kb VtC XES for NiR active site with side-on Cu(I)-(NO+), end-on Cu(I)-(NO+), and Cu(II)-NO2. Structures of the models are presented on the right. A broadening of 2.5 eV was applied, but no energy calibration was conducted. Reprinted with permission from Phu, P. N.; Gutierrez, C. E.; Kundu, S.; Sokaras, D.; Kroll, T.; Warren, T. H.; Stieber, S. C. E. Quantification of Ni–N–O Bond Angles and NO Activation by X-Ray Emission Spectroscopy. Inorg. Chem. 2021, 60(2), 736–744. Copyright 2021 American Chemical Society.

Computational Methods in Organometallic Chemistry

199

the model. Truncation is becoming less justifiable, as the full molecule can typically be used for calculations given current accessibility of computational power. The second common error in mechanistic studies is the improper selection of basis sets and functionals (see Sections 1.07.2.1 and 1.07.2.2). Using a simple functional and double-z basis set (such as BP86 and 6-31G ) may result in a faster calculation, but could result in problematically low accuracy. For energy calculations, B3LYP is commonly used and ideally a triple-z basis set is used for at least the metal and primary atoms of interest. However, particularly in high-spin systems, B3LYP can have large errors. As always, the more experimental benchmarks that can be used to validate the computational work, the more chemically reasonable the result. For example, if a reaction was monitored by infrared spectroscopy, IR calculations could provide an additional experimental comparison to observed intermediates in a proposed mechanism. If calculated trends are in agreement with experimental data, this could lend support to the computational model. Some of the additional factors to consider include quantum chemical methodology, entropic/thermal effects, the kinetic model, solvation, conformational complexity, the microscopic model, selectivity and accuracy, electron transfer, and dynamics, to name a few.293 A few will be highlighted here. Experimentalists are well aware of changing temperature to speed up or slow down a reaction, and so it is not surprising that temperature can also significantly affect computational thermochemical output, which are typically output at STP by default. Therefore, it is crucial to consider translational entropy corrections, which range from −15 to −25 kcal mol−1 from 100  C to 300  C for example.292 Similarly, when multiple molecular fragments are combined, there are entropic penalties of approximately 11 kcal mol−1 to combine 2 fragments, approximately 22 kcal mol−1 to combine 3 fragments, and approximately 33 kcal mol−1 to combine 4 fragments.292 This can greatly affect the calculated Gibbs free energy of transition states for a reaction of interest. The solvent is also a key component of experimental systems that is challenging to computationally model. Since solvent-solute interactions are weak, the solvent can only be explicitly modeled if doing molecular dynamics calculations which are very computationally expensive. Generally, DFT calculations are conducted with a polarizable continuum model (PCM) of the solvent, which allows one to choose a dielectric constant that matches the experimental solvent.294,295 COSMO and CPCM are common models used for DFT of organometallic systems.47 Perhaps the biggest challenge is verifying that an optimized geometry is in a global versus local energy minimum, and this is still a very active area of research.296,297 This is particularly difficult in systems with bulky ligands, which offer barriers to rapid rotation and many local minima. Practically speaking, the main strategy is to empirically compare results with known data. As much as possible, proposed structures should be compared to known structures with crystallographic data available to compare similar bond lengths and angles to see if the calculated ones are chemically reasonable. The electronic structure can also be used as an indicator to evaluate how significant a change such as a rotation or other small change is. Additionally, using fundamental organometallic knowledge of ligand field theory and electron counting can aid in evaluating how reasonable a result is. If one has the time and computational resources, molecular dynamics or Monte Carlo methods may allow for more possible structures to be explicitly probed allowing for higher certainty that the local and global minima calculated are actually the lowest energy states.

1.07.6.1

Computational methods for calculating reaction mechanisms

Mechanisms can be particularly challenging to computationally predict, partially due to the need to account for static and dynamic correlation energies for calculating barrier heights in reactions.293,298,299 Multireference methods such as MCSCF, CASPT2, MRMP2 and MRCI offer the most accurate results, however they are so computationally expensive and time consuming that they are limited to single point calculations.108,110,112,300–310 Additionally, the geometries for transition states are found through DFT first, which potentially biases the calculation to inaccurate transition states. Recent work suggests that MC-PDF has similar accuracy and is less sensitive than CSPT2 to the choice of basis set.311 There is some use of molecular dynamics (MD) for calculating organometallic mechanisms, such as for calculating the energy for inner sphere reorganization in the cobalt hexamine(II/III) self exchange reaction.312

1.07.6.1.1

Case study 13: Mechanism of CdCO2 bond formation at Cu, Rh and Pd

CO2 activation is an area of considerable interest for organometallic chemistry, however mechanisms are poorly understood which makes the rational design of catalysts challenging. For CdCO2 bond formation, both inner sphere and outer sphere mechanisms have been experimentally observed, but the origins are unclear. An inner sphere mechanism involves the metal during the CdCO2 bond formation step, while the outer sphere mechanism only involves the two carbon atoms present in the resulting CdCO2 bond (Fig. 18). Recent computational work using DFT with an IEFPCM implicit solvent model investigated the inner versus outer sphere mechanistic pathways for Cu, Rh and Pd with N-heterocyclic carbene (NHC), phosphine, and pincer ligands.313 Ground state and transition state geometry optimizations were conducted with B3LYP as the functional,54,165 SDD basis set on the metals (recall that these use ECP or effective core potentials),166,200 and 6-31 ++G(d,p) on the other atoms.174,314–317 It was crucial to choose the SDD basis set with the effective core potential even for Cu (although an all-electron basis set would have been more affordable), because then the same basis set can be used as heavier elements such as Rh and Pd. Additionally a dispersion correction D3,318,319 and solvent model IEFPCM were used.320 When comparing energies, it is critical that the same number of atoms are present in the systems being compared, so free energy barriers (DG) were referenced to a starting point of free CO2 and the metal alkyl. Results suggest that the outer sphere mechanism is favored for benzylic sp3 carbon nucleophiles with Pd and Rh due to lower steric pressure (Fig. 18). By contrast, for Cu with benzylic sp3 carbon nucleophiles, the outer sphere mechanism is preferred for bulky ligands, while inner sphere is favored with smaller ligands. For sp2 carbon nucleophiles, the inner sphere mechanism is favored for

200

Computational Methods in Organometallic Chemistry

Fig. 18 Mechanism for inner-sphere (left) and outer-sphere (right) CdCO2 bond formation (top) and geometry optimized transition states for CdCO2 bond formation in Pd complexes (bottom). Figure used with permission from García-López, D.; Pavlovic, L.; Hopmann, K. H. To Bind or Not to Bind: Mechanistic Insights into C–CO2 Bond Formation with Late Transition Metals. Organometallics 2020, 39(8), 1339–1347, 10.1021/acs.organomet.0c00090.

Cu, Rh and Pd. Remaining areas for study include the broad applicability of these findings for other transition metal centers and substrates. Based on the current study alone, it is unclear if these results only apply to the specific ligand systems studied.

1.07.7

Current limitations and outlook

Active areas of development in computational organometallic chemistry include methods for more accurately calculating electronic spectroscopy,113 multireference methods for f-elements,91 multicomponent mechanisms,293,321,322 and improving predictive computational capacity,323 for example. While it may be tempting for a computational novice to embark on calculations alone, it is critical to remember that the computer will produce an answer, whether it is chemically reasonable or not! Calculations and results are highly dependent on functionals and basis sets, so it is important to use basis sets and functionals designed for or that are well-tested for the specific application probed (for example geometry optimization versus EPR calculation). When determining electronic structures, it is best if a computational model can be additionally benchmarked by calculating spectroscopic parameters and comparing with experimental data such as EPR, Mössbauer, XAS, IR, or UV–vis. It is still quite difficult to computationally predict a reaction or reaction mechanism and be sure that real global or representative local minima in systems with complicated ligand spheres are being probed, so it is critical to benchmark with as much experimental data as possible. When problems are encountered with DFT, ab initio or molecular dynamics approaches may be necessary, but it is important to balance the time investment versus what will be learned. Overall, DFT still has many applications for organometallic chemists hoping to gain a better understanding of a system and the overall accessibility of computational methods and applications are likely to keep improving.

1.07.7.1

Case study 14: The ever elusive Grignard reaction

The reader is left with one final case study, which illustrates some of the aforementioned existing challenges in computational chemistry as it relates to organometallic systems. One of the first organometallic reactions that students learn is the Grignard reaction which was discovered in 1900,324 so it is in some ways surprising that the exact mechanism is still not fully understood. Recent computational work has made significant headway in understanding the various factors involved.322 Typically, the Grignard reagent is thought of as a nucleophile (polar mechanism), however electron-transfer (radical mechanism) may also contribute (Fig. 19),325 which would affect rational design and variations of the reaction. Ab initio molecular dynamics (AIMD) calculations demonstrated the ability for Mg to have a range of solvent coordination numbers in the first coordination sphere that is dependent on chloride versus methyl coordination.326 In the current case study, the addition of enhanced-sampling methods allowed for a broader range of solvent dynamics and solvation to be probed.327

Computational Methods in Organometallic Chemistry

201

Fig. 19 Possible mechanisms for the Grignard reaction, including the polar (top) and radical (bottom) mechanisms.

This study highlights forefront work in combining ab initio computational methods with molecular dynamics calculations, which allow for movements of molecules to be probed. Structures were determined by DFT, then AIMD calculations were conducted within a space of 25.2  15.0  15.0 A˚ 3 that contained one Mg center (as in Fig. 19) and 41 THF molecules to account for solvent. This allowed for possible solvent participation in the reaction to be more accurately probed than in usual DFT calculations. Several possible pathways and transition states were probed for polar and radical Grignard mechanisms. Similar activation energies found by this study for differing transition states suggest that there may be multiple parallel reaction paths in solution. Results suggest that the polar mechanism is highly dependent on the solvent, justifying the use of the explicit solvent model and AIMD, and that the reaction is slowed by more bulky substituents. The radical mechanism is also dependent on the solvent, but is favored for substrates with empty p-orbitals that have low energies and bulky substituents. Overall, the more solvated a Mg species is, the more reactive it was found to be. Without the ability for AIMD to probe multiple explicit solvent molecules, this would not have been possible to determine with DFT alone. However, the effects of the individual simultaneous reactions and their effects on the overall reaction are still under investigation.

Acknowledgments S.C.E.S. acknowledges financial support from NSF CAREER (CHE-1847926), CSU ARI (21-04-108) and NSF XSEDE (CHE160059, ACI-1548562).

References 1. Cramer, C. J. Essentials of Computational Chemistry: Theories and Models, 2nd ed.; Wiley-VCH Verlag GmbH & Co. KGaA: New York, 2004. 2. Hopmann, K. H. How To Make Your Computational Paper Interesting and Have It Published. Organometallics 2019, 38 (3), 603–605. https://doi.org/10.1021/acs. organomet.8b00942. 3. Wiest, O., Wu, Y., Eds.; In Computational Organometallic Chemistry; Springer-Verlag: Berlin, Heidelberg, 2012https://doi.org/10.1007/978-3-642-25258-7. 4. Macgregor, S. A., Eisenstein, O., Eds.; In Computational Studies in Organometallic Chemistry; Structure and Bonding Springer International Publishing, 2016https://doi.org/ 10.1007/978-3-319-31638-3. 5. Cundari, T. R. Computational Organometallic Chemistry; CRC Press, 2001. 6. Chermette, H. Density Functional Theory: A Powerful Tool for Theoretical Studies in Coordination Chemistry. Coord. Chem. Rev. 1998, 178–180, 699–721. https://doi.org/ 10.1016/S0010-8545(98)00179-9. 7. Lovell, T.; Himo, F.; Han, W.-G.; Noodleman, L. Density Functional Methods Applied to Metalloenzymes. Coord. Chem. Rev. 2003, 238–239, 211–232. https://doi.org/ 10.1016/S0010-8545(02)00331-4. 8. Ciofini, I.; Daul, C. A. DFT Calculations of Molecular Magnetic Properties of Coordination Compounds. Coord. Chem. Rev. 2003, 238–239, 187–209. https://doi.org/10.1016/ S0010-8545(02)00330-2. 9. Autschbach, J.; Ziegler, T. Double Perturbation Theory: A Powerful Tool in Computational Coordination Chemistry. Coord. Chem. Rev. 2003, 238–239, 83–126. https://doi.org/ 10.1016/S0010-8545(02)00287-4. 10. Bickelhaupt, F. M.; Baerends, E. J. Kohn-Sham Density Functional Theory: Predicting and Understanding Chemistry. In Reviews in Computational Chemistry, John Wiley & Sons, Ltd, 2007;; pp 1–86. https://doi.org/10.1002/9780470125922.ch1. 11. Yu, H. S.; Li, S. L.; Truhlar, D. G. Perspective: Kohn-Sham Density Functional Theory Descending a Staircase. J. Chem. Phys. 2016, 145 (13), 130901. https://doi.org/ 10.1063/1.4963168.

202

Computational Methods in Organometallic Chemistry

12. Runge, E.; Gross, E. K. U. Density-Functional Theory for Time-Dependent Systems. Phys. Rev. Lett. 1984, 52 (12), 997–1000. https://doi.org/10.1103/PhysRevLett.52.997. 13. Casida, M. E. Time-Dependent Density Functional Response Theory for Molecules. In Recent Advances in Density Functional Methods, Recent Advances in Computational Chemistry World Scientific, 1995; vol. 1; pp 155–192. https://doi.org/10.1142/9789812830586_0005. 14. Furche, F. On the Density Matrix Based Approach to Time-Dependent Density Functional Response Theory. J. Chem. Phys. 2001, 114 (14), 5982–5992. https://doi.org/ 10.1063/1.1353585. 15. Bauernschmitt, R.; Ahlrichs, R. Treatment of Electronic Excitations Within the Adiabatic Approximation of Time Dependent Density Functional Theory. Chem. Phys. Lett. 1996, 256 (4–5), 454–464. https://doi.org/10.1016/0009-2614(96)00440-X. 16. Bauernschmitt, R.; Häser, M.; Treutler, O.; Ahlrichs, R. Calculation of Excitation Energies Within Time-Dependent Density Functional Theory Using Auxiliary Basis Set Expansions. Chem. Phys. Lett. 1997, 264 (6), 573–578. https://doi.org/10.1016/S0009-2614(96)01343-7. 17. Dreuw, A.; Head-Gordon, M. Single-Reference Ab Initio Methods for the Calculation of Excited States of Large Molecules. Chem. Rev. 2005, 105 (11), 4009–4037. https://doi. org/10.1021/cr0505627. 18. Grimme, S.; Parac, M. Substantial Errors from Time-Dependent Density Functional Theory for the Calculation of Excited States of Large p Systems. ChemPhysChem 2003, 4 (3), 292–295. https://doi.org/10.1002/cphc.200390047. 19. Kohn, W.; Sham, L. J. Self-Consistent Equations Including Exchange and Correlation Effects. Phys. Rev. 1965, 140 (4A), A1133–A1138. https://doi.org/10.1103/PhysRev.140. A1133. 20. Mardirossian, N.; Head-Gordon, M. Thirty Years of Density Functional Theory in Computational Chemistry: An Overview and Extensive Assessment of 200 Density Functionals. Mol. Phys. 2017, 115 (19), 2315–2372. https://doi.org/10.1080/00268976.2017.1333644. 21. Sousa, S. F.; Fernandes, P. A.; Ramos, M. J. General Performance of Density Functionals. J. Phys. Chem. A 2007, 111 (42), 10439–10452. https://doi.org/10.1021/ jp0734474. 22. Rappoport, D.; Crawford, N. R. M.; Furche, F.; Burke, K. Approximate Density Functionals: Which Should I Choose?In Encyclopedia of Inorganic Chemistry, American Cancer Society, 2009https://doi.org/10.1002/0470862106.ia615. 23. Vosko, S. H.; Wilk, L.; Nusair, M. Accurate Spin-Dependent Electron Liquid Correlation Energies for Local Spin Density Calculations: A Critical Analysis. Can. J. Phys. 2011https://doi.org/10.1139/p80-159. 24. Perdew, J. P.; Wang, Y. Accurate and Simple Analytic Representation of the Electron-Gas Correlation Energy. Phys. Rev. B 1992, 45 (23), 13244–13249. https://doi.org/ 10.1103/PhysRevB.45.13244. 25. Burke, K.; Perdew, J. P.; Wang, Y. Derivation of a Generalized Gradient Approximation: The PW91 Density Functional. In Electronic Density Functional Theory: Recent Progress and New Directions; Dobson, J. F., Vignale, G., Das, M. P., Eds.; Springer US: Boston, MA, 1998;; pp 81–111. https://doi.org/10.1007/978-1-4899-0316-7_7. 26. Tao, J.; Perdew, J. P.; Staroverov, V. N.; Scuseria, G. E. Climbing the Density Functional Ladder: Nonempirical Meta—Generalized Gradient Approximation Designed for Molecules and Solids. Phys. Rev. Lett. 2003, 91 (14), 146401. https://doi.org/10.1103/PhysRevLett.91.146401. 27. Neese, F.; Solomon, E. I. Interpretation and Calculation of Spin-Hamiltonian Parameters in Transition Metal Complexes. In Magnetism: Molecules to Materials IV, John Wiley & Sons, Ltd, 2003;; pp 345–466. https://doi.org/10.1002/3527600698.ch9. 28. Jensen, K. P. Bioinorganic Chemistry Modeled With the TPSSh Density Functional. Inorg. Chem. 2008, 47 (22), 10357–10365. https://doi.org/10.1021/ic800841t. 29. Paier, J.; Marsman, M.; Kresse, G. Why Does the B3LYP Hybrid Functional Fail for Metals?J. Chem. Phys. 2007, 127 (2), 024103https://doi.org/10.1063/1.2747249. 30. Salomon, O.; Reiher, M.; Hess, B. A. Assertion and Validation of the Performance of the B3LYP ? Functional for the First Transition Metal Row and the G2 Test Set. J. Chem. Phys. 2002, 117 (10), 4729–4737. https://doi.org/10.1063/1.1493179. 31. Adamo, C.; Barone, V. Toward Reliable Density Functional Methods Without Adjustable Parameters: The PBE0 Model. J. Chem. Phys. 1999, 110 (13), 6158–6170. https://doi. org/10.1063/1.478522. 32. Slater, J. C. Atomic Shielding Constants. Phys. Rev. 1930, 36 (1), 57–64. https://doi.org/10.1103/PhysRev.36.57. 33. Gill, P. M. W. Molecular Integrals Over Gaussian Basis Functions. In Advances in Quantum Chemistry; Sabin, J. R., Zerner, M. C., Eds.; 25; Academic Press, 1994; pp 141–205. https://doi.org/10.1016/S0065-3276(08)60019-2. 34. Weigend, F.; Ahlrichs, R. Balanced Basis Sets of Split Valence, Triple Zeta Valence and Quadruple Zeta Valence Quality for H to Rn: Design and Assessment of Accuracy. Phys. Chem. Chem. Phys. 2005, 7 (18), 3297–3305. https://doi.org/10.1039/B508541A. 35. Pantazis, D. A.; Chen, X.-Y.; Landis, C. R.; Neese, F. All-Electron Scalar Relativistic Basis Sets for Third-Row Transition Metal Atoms. J. Chem. Theory Comput. 2008, 4 (6), 908–919. https://doi.org/10.1021/ct800047t. 36. Johnson, B. G.; Gill, P. M. W.; Pople, J. A. The Performance of a Family of Density Functional Methods. J. Chem. Phys. 1993, 98 (7), 5612–5626. https://doi.org/ 10.1063/1.464906. 37. Johnson, B. G.; Gill, P. M. W.; Pople, J. A. Preliminary Results on the Performance of a Family of Density Functional Methods. J. Chem. Phys. 1992, 97 (10), 7846–7848. https://doi.org/10.1063/1.463975. 38. Küchle, W.; Dolg, M.; Stoll, H.; Preuss, H. Ab Initio Pseudopotentials for Hg through Rn. Mol. Phys. 1991, 74 (6), 1245–1263. https://doi.org/10.1080/00268979100102941. 39. Küchle, W.; Dolg, M.; Stoll, H.; Preuss, H. Energy-adjusted Pseudopotentials for the Actinides. Parameter Sets and Test Calculations for Thorium and Thorium Monoxide. J. Chem. Phys. 1994, 100 (10), 7535–7542. https://doi.org/10.1063/1.466847. 40. Petersson, G. A.; Bennett, A.; Tensfeldt, T. G.; Al-Laham, M. A.; Shirley, W. A.; Mantzaris, J. A Complete Basis Set Model Chemistry. I. The Total Energies of Closed-shell Atoms and Hydrides of the First-Row Elements. J. Chem. Phys. 1988, 89 (4), 2193–2218. https://doi.org/10.1063/1.455064. 41. Hay, P. J.; Martin, R. L. Theoretical Studies of the Structures and Vibrational Frequencies of Actinide Compounds Using Relativistic Effective Core Potentials with Hartree–Fock and Density Functional Methods: UF6, NpF6, and PuF6. J. Chem. Phys. 1998, 109 (10), 3875–3881. https://doi.org/10.1063/1.476988. 42. Cao, X.; Dolg, M.; Stoll, H. Valence Basis Sets for Relativistic Energy-Consistent Small-Core Actinide Pseudopotentials. J. Chem. Phys. 2002, 118 (2), 487–496. https://doi.org/ 10.1063/1.1521431. 43. Dolg, M.; Cao, X. Accurate Relativistic Small-Core Pseudopotentials for Actinides. Energy Adjustment for Uranium and First Applications to Uranium Hydride. J. Phys. Chem. A 2009, 113 (45), 12573–12581. https://doi.org/10.1021/jp9044594. 44. Pantazis, D. A.; Neese, F. All-Electron Scalar Relativistic Basis Sets for the Actinides. J. Chem. Theory Comput. 2011, 7 (3), 677–684. https://doi.org/10.1021/ct100736b. 45. Kossmann, S.; Neese, F. Efficient Structure Optimization with Second-Order Many-Body Perturbation Theory: The RIJCOSX-MP2 Method. J. Chem. Theory Comput. 2010, 6 (8), 2325–2338. https://doi.org/10.1021/ct100199k. 46. Izsák, R.; Hansen, A.; Neese, F. The Resolution of Identity and Chain of Spheres Approximations for the LPNO-CCSD Singles Fock Term. Mol. Phys. 2012, 110 (19–20), 2413–2417. https://doi.org/10.1080/00268976.2012.687466. 47. Klamt, A.; Schüürmann, G. COSMO: A New Approach to Dielectric Screening in Solvents with Explicit Expressions for the Screening Energy and Its Gradient. J. Chem. Soc., Perkin Trans. 2 1993, (5), ;799–805. https://doi.org/10.1039/P29930000799. 48. Grimme, S. Density Functional Theory With London Dispersion Corrections. WIREs Comput. Mol. Sci. 2011, 1 (2), 211–228. https://doi.org/10.1002/wcms.30. 49. Evans, D. F. 400. The Determination of the Paramagnetic Susceptibility of Substances in Solution by Nuclear Magnetic Resonance. J. Chem. Soc. 1959, ;2003–2005. https:// doi.org/10.1039/JR9590002003. No. 0. 50. Piguet, C. Paramagnetic Susceptibility by NMR: The “Solvent Correction” Removed for Large Paramagnetic Molecules. J. Chem. Educ. 1997, 74 (7), 815. https://doi.org/ 10.1021/ed074p815. 51. Massa, W. Crystal Structure Determination, 2nd ed.; Springer-Verlag: Berlin Heidelberg, 2004https://doi.org/10.1007/978-3-662-06431-3.

Computational Methods in Organometallic Chemistry

203

52. Perdew, J. P. Density-Functional Approximation for the Correlation Energy of the Inhomogeneous Electron Gas. Phys. Rev. B 1986, 33 (12), 8822–8824. https://doi.org/ 10.1103/PhysRevB.33.8822. 53. Perdew, J.; Erratum, P. Density-Functional Approximation for the Correlation Energy of the Inhomogeneous Electron Gas. Phys. Rev. B 1986, 34 (10), 7406. https://doi.org/ 10.1103/PhysRevB.34.7406. 54. Lee, C.; Yang, W.; Parr, R. G. Development of the Colle-Salvetti Correlation-Energy Formula into a Functional of the Electron Density. Phys. Rev. B 1988, 37 (2), 785–789. https://doi.org/10.1103/PhysRevB.37.785. 55. Schäfer, A.; Huber, C.; Ahlrichs, R. Fully Optimized Contracted Gaussian Basis Sets of Triple Zeta Valence Quality for Atoms Li to Kr. J. Chem. Phys. 1994, 100 (8), 5829–5835. https://doi.org/10.1063/1.467146. 56. Schäfer, A.; Horn, H.; Ahlrichs, R. Fully Optimized Contracted Gaussian Basis Sets for Atoms Li to Kr. J. Chem. Phys. 1992, 97 (4), 2571–2577. https://doi.org/ 10.1063/1.463096. 57. Eichkorn, K.; Weigend, F.; Treutler, O.; Ahlrichs, R. Auxiliary Basis Sets for Main Row Atoms and Transition Metals and Their Use to Approximate Coulomb Potentials. Theor. Chem. Acc. 1997, 97 (1), 119–124. https://doi.org/10.1007/s002140050244. 58. ORCA Input Library, https://sites.google.com/site/orcainputlibrary/home. accessed 2021 -05 -26. 59. Neese, F. Orca: An Ab Initio, DFT and Semiempirical Electronic Structure Package; Max Planck Institute for Chemical Energy Conversion: Mülheim an der Ruhr, Germany, 2021. 60. Giannozzi, P.; Baroni, S.; Bonini, N.; Calandra, M.; Car, R.; Cavazzoni, C.; Ceresoli, D.; Chiarotti, G. L.; Cococcioni, M.; Dabo, I.; Corso, A. D.; de Gironcoli, S.; Fabris, S.; Fratesi, G.; Gebauer, R.; Gerstmann, U.; Gougoussis, C.; Kokalj, A.; Lazzeri, M.; Martin-Samos, L.; Marzari, N.; Mauri, F.; Mazzarello, R.; Paolini, S.; Pasquarello, A.; Paulatto, L.; Sbraccia, C.; Scandolo, S.; Sclauzero, G.; Seitsonen, A. P.; Smogunov, A.; Umari, P.; Wentzcovitch, R. M. QUANTUM ESPRESSO: A Modular and Open-Source Software Project for Quantum Simulations of Materials. J. Phys. Condens. Matter 2009, 21 (39), 395502. https://doi.org/10.1088/0953-8984/21/39/395502. 61. Barca, G. M. J.; Bertoni, C.; Carrington, L.; Datta, D.; De Silva, N.; Deustua, J. E.; Fedorov, D. G.; Gour, J. R.; Gunina, A. O.; Guidez, E.; Harville, T.; Irle, S.; Ivanic, J.; Kowalski, K.; Leang, S. S.; Li, H.; Li, W.; Lutz, J. J.; Magoulas, I.; Mato, J.; Mironov, V.; Nakata, H.; Pham, B. Q.; Piecuch, P.; Poole, D.; Pruitt, S. R.; Rendell, A. P.; Roskop, L. B.; Ruedenberg, K.; Sattasathuchana, T.; Schmidt, M. W.; Shen, J.; Slipchenko, L.; Sosonkina, M.; Sundriyal, V.; Tiwari, A.; Galvez Vallejo, J. L.; Westheimer, B.; Włoch, M.; Xu, P.; Zahariev, F.; Gordon, M. S. Recent Developments in the General Atomic and Molecular Electronic Structure System. J. Chem. Phys. 2020, 152 (15), 154102. https://doi.org/10.1063/5.0005188. 62. Pettersen, E. F.; Goddard, T. D.; Huang, C. C.; Couch, G. S.; Greenblatt, D. M.; Meng, E. C.; Ferrin, T. E. Chimera. J. Comput. Chem. 2004, 13, 1605. 63. Hanwell, M. D.; Curtis, D. E.; Lonie, D. C.; Vandermeersch, T.; Zurek, E.; Hutchison, G. R. Avogadro: An Advanced Semantic Chemical Editor, Visualization, and Analysis Platform. J. Cheminformatics 2012, 4 (1), 17. https://doi.org/10.1186/1758-2946-4-17. 64. Humphrey, W.; Dalke, A.; Schulten, K. VMD—Visual Molecular Dynamics. J. Mol. Graph. 1996, 14, 33–38. 65. Varshney, A.; Brooks, F. P.; Wright, W. V. Linearly Scalable Computation of Smooth Molecular Surfaces. IEEE Comput. Graph. Appl. 1994, 14, 19–25. 66. Frishman, D.; Argos, P. Knowledge-Based Secondary Structure Assignment. Proteins Struct. Funct. Genet. 1995, 23, 566–579. 67. Sanner, M.; Olsen, A.; Spehner, J.-C. Fast and Robust Computation of Molecular Surfaces. In Proceedings of the 11th ACM Symposium on Computational Geometry, ACM: New York, 1995;; pp C6–C7. 68. Stone, J. An Efficient Library for Parallel Ray Tracing and Animation; Master’s Thesis Computer Science Department, University of Missouri-Rolla, 1998. 69. Sharma, R.; Zeller, M.; Pavlovic, V. I.; Huang, T. S.; Lo, Z.; Chu, S.; Zhao, Y.; Phillips, J. C.; Schulten, K. Speech/Gesture Interface to a Visual-Computing Environment. IEEECGA 2000, 20, 29–37. 70. Stone, J.; Gullingsrud, J.; Grayson, P.; Schulten, K. A System for Interactive Molecular Dynamics Simulation; In ACM Symposium on Interactive 3D Graphics; Hughes, J. F., Séquin, C. H., Eds.; 2001, ACM SIGGRAPH: New York, 2001; pp 191–194. 71. Eargle, J.; Wright, D.; Luthey-Schulten, Z. Multiple Alignment of Protein Structures and Sequences for VMD. Bioinformatics 2006, 22 (4), 504–506. 72. Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Scalmani, G.; Barone, V.; Petersson, G. A.; Nakatsuji, H.; Li, X.; Caricato, M.; Marenich, A. V.; Bloino, J.; Janesko, B. G.; Gomperts, R.; Mennucci, B.; Hratchian, H. P.; Ortiz, J. V.; Izmaylov, A. F.; Sonnenberg, J. L.; Williams, ; Ding, F.; Lipparini, F.; Egidi, F.; Goings, J.; Peng, B.; Petrone, A.; Henderson, T.; Ranasinghe, D.; Zakrzewski, V. G.; Gao, J.; Rega, N.; Zheng, G.; Liang, W.; Hada, M.; Ehara, M.; Toyota, K.; Fukuda, R.; Hasegawa, J.; Ishida, M.; Nakajima, T.; Honda, Y.; Kitao, O.; Nakai, H.; Vreven, T.; Throssell, K.; Montgomery, J. A., Jr.; Peralta, J. E.; Ogliaro, F.; Bearpark, M. J.; Heyd, J. J.; Brothers, E. N.; Kudin, K. N.; Staroverov, V. N.; Keith, T. A.; Kobayashi, R.; Normand, J.; Raghavachari, K.; Rendell, A. P.; Burant, J. C.; Iyengar, S. S.; Tomasi, J.; Cossi, M.; Millam, J. M.; Klene, M.; Adamo, C.; Cammi, R.; Ochterski, J. W.; Martin, R. L.; Morokuma, K.; Farkas, O.; Foresman, J. B.; Fox, D. J. Gaussian 16 Rev. C.01; Wallingford, CT https://gaussian.com/g03citation/. 73. Werner, H.-J.; Knowles, P. J.; Knizia, G.; Manby, F. R.; Schütz, M. Molpro: A General-Purpose Quantum Chemistry Program Package. WIREs Comput. Mol. Sci. 2012, 2, 242–253. 74. te Velde, G.; Bickelhaupt, F. M.; Baerends, E. J.; Fonseca Guerra, C.; van Gisbergen, S. J. A.; Snijders, J. G.; Ziegler, T. Chemistry With ADF. J. Comput. Chem. 2001, 22 (9), 931–967. https://doi.org/10.1002/jcc.1056. 75. Noodleman, L.; Peng, C. Y.; Case, D. A.; Mouesca, J.-M. Orbital Interactions, Electron Delocalization and Spin Coupling in Iron-Sulfur Clusters. Coord. Chem. Rev. 1995, 144, 199–244. https://doi.org/10.1016/0010-8545(95)07011-L. 76. Ginsberg, A. P. Magnetic Exchange in Transition Metal Complexes. 12. Calculation of Cluster Exchange Coupling Constants with the X.Alpha.-Scattered Wave Method. J. Am. Chem. Soc. 1980, 102 (1), 111–117. https://doi.org/10.1021/ja00521a020. 77. Kirchner, B.; Wennmohs, F.; Ye, S.; Neese, F. Theoretical Bioinorganic Chemistry: The Electronic Structure Makes a Difference. Curr. Opin. Chem. Biol. 2007, 11 (2), 134–141. https://doi.org/10.1016/j.cbpa.2007.02.026. 78. Soda, T.; Kitagawa, Y.; Onishi, T.; Takano, Y.; Shigeta, Y.; Nagao, H.; Yoshioka, Y.; Yamaguchi, K. Ab Initio Computations of Effective Exchange Integrals for H–H, H–He–H and Mn2O2 Complex: Comparison of Broken-Symmetry Approaches. Chem. Phys. Lett. 2000, 319 (3), 223–230. https://doi.org/10.1016/S0009-2614(00)00166-4. 79. Stieber, S. C. E.; Milsmann, C.; Hoyt, J. M.; Turner, Z. R.; Finkelstein, K. D.; Wieghardt, K.; DeBeer, S.; Chirik, P. J. Bis(Imino)Pyridine Iron Dinitrogen Compounds Revisited: Differences in Electronic Structure Between Four- and Five-Coordinate Derivatives. Inorg. Chem. 2012, 51 (6), 3770–3785. https://doi.org/10.1021/ic202750n. 80. Tondreau, A. M.; Stieber, S. C. E.; Milsmann, C.; Lobkovsky, E.; Weyhermüller, T.; Semproni, S. P.; Chirik, P. J. Oxidation and Reduction of Bis(Imino)Pyridine Iron Dinitrogen Complexes: Evidence for Formation of a Chelate Trianion. Inorg. Chem. 2013, 52 (2), 635–646. https://doi.org/10.1021/ic301675t. 81. Milsmann, C.; Turner, Z. R.; Semproni, S. P.; Chirik, P. J. Azo N ¼ N Bond Cleavage with a Redox-Active Vanadium Compound Involving Metal–Ligand Cooperativity. Angew. Chem. Int. Ed. 2012, 51 (22), 5386–5390. https://doi.org/10.1002/anie.201201085. 82. Darmon, J. M.; Stieber, S. C. E.; Sylvester, K. T.; Fernández, I.; Lobkovsky, E.; Semproni, S. P.; Bill, E.; Wieghardt, K.; DeBeer, S.; Chirik, P. J. Oxidative Addition of Carbon–Carbon Bonds with a Redox-Active Bis(Imino)Pyridine Iron Complex. J. Am. Chem. Soc. 2012, 134 (41), 17125–17137. https://doi.org/10.1021/ja306526d. 83. Yu, R. P.; Darmon, J. M.; Milsmann, C.; Margulieux, G. W.; Stieber, S. C. E.; DeBeer, S.; Chirik, P. J. Catalytic Hydrogenation Activity and Electronic Structure Determination of Bis(Arylimidazol-2-Ylidene)Pyridine Cobalt Alkyl and Hydride Complexes. J. Am. Chem. Soc. 2013, 135 (35), 13168–13184. https://doi.org/10.1021/ja406608u. 84. Peterson, P. O.; Rummelt, S. M.; Wile, B. M.; Stieber, S. C. E.; Zhong, H.; Chirik, P. J. Direct Observation of Transmetalation from a Neutral Boronate Ester to a Pyridine(Diimine) Iron Alkoxide. Organometallics 2020, 39 (1), 201–205. https://doi.org/10.1021/acs.organomet.9b00733. 85. Darmon, J. M.; Yu, R. P.; Semproni, S. P.; Turner, Z. R.; Stieber, S. C. E.; DeBeer, S.; Chirik, P. J. Electronic Structure Determination of Pyridine N-Heterocyclic Carbene Iron Dinitrogen Complexes and Neutral Ligand Derivatives. Organometallics 2014, 33 (19), 5423–5433. https://doi.org/10.1021/om500727t. 86. Ray, K.; Weyhermüller, T.; Neese, F.; Wieghardt, K. Electronic Structure of Square Planar Bis(Benzene-1,2-Dithiolato)Metal Complexes [M(L)2]z (z ¼ 2−, 1−, 0; M ¼ Ni, Pd, Pt, Cu, Au): An Experimental, Density Functional, and Correlated Ab Initio Study. Inorg. Chem. 2005, 44 (15), 5345–5360. https://doi.org/10.1021/ic0507565.

204

Computational Methods in Organometallic Chemistry

87. Kundu, S.; Phu, P. N.; Ghosh, P.; Kozimor, S. A.; Bertke, J. A.; Stieber, S. C. E.; Warren, T. H. Nitrosyl Linkage Isomers: NO Coupling to N 2 O at a Mononuclear Site. J. Am. Chem. Soc. 2019, 141 (4), 1415–1419. https://doi.org/10.1021/jacs.8b09769. 88. Kundu, S.; Stieber, S. C. E.; Ferrier, M. G.; Kozimor, S. A.; Bertke, J. A.; Warren, T. H. Redox Non-Innocence of Nitrosobenzene at Nickel. Angew. Chem. Int. Ed. 2016, 55 (35), 10321–10325. https://doi.org/10.1002/anie.201605026. 89. Meskaldji, S.; Zaiter, A.; Belkhiri, L.; Boucekkine, A. A Relativistic DFT Study of Magnetic Exchange Coupling in Ketimide Bimetallic Uranium(IV) Complexes. Theor. Chem. Accounts 2012, 131 (3), 1151. https://doi.org/10.1007/s00214-012-1151-9. 90. Reger, D. L.; Pascui, A. E.; Smith, M. D.; Jezierska, J.; Ozarowski, A. Halide and Hydroxide Linearly Bridged Bimetallic Copper(II) Complexes: Trends in Strong Antiferromagnetic Superexchange Interactions. Inorg. Chem. 2012, 51 (15), 7966–7968. https://doi.org/10.1021/ic301321r. 91. Kaltsoyannis, N. Transuranic Computational Chemistry. Chem. Eur. J. 2018, 24 (12), 2815–2825. https://doi.org/10.1002/chem.201704445. 92. Roy, L. E.; Hay, P. J.; Martin, R. L. Revised Basis Sets for the LANL Effective Core Potentials. J. Chem. Theory Comput. 2008, 4 (7), 1029–1031. https://doi.org/10.1021/ ct8000409. 93. Hay, P. J.; Wadt, W. R. Ab Initio Effective Core Potentials for Molecular Calculations. Potentials for the Transition Metal Atoms Sc to Hg. J. Chem. Phys. 1985, 82 (1), 270–283. https://doi.org/10.1063/1.448799. 94. Martin, J. M. L.; Sundermann, A. Correlation Consistent Valence Basis Sets for Use With the Stuttgart–Dresden–Bonn Relativistic Effective Core Potentials: The Atoms Ga–Kr and In–Xe. J. Chem. Phys. 2001, 114 (8), 3408–3420. https://doi.org/10.1063/1.1337864. 95. Fuentealba, P.; Preuss, H.; Stoll, H.; Von Szentpály, L. A Proper Account of Core-Polarization with Pseudopotentials: Single Valence-Electron Alkali Compounds. Chem. Phys. Lett. 1982, 89 (5), 418–422. https://doi.org/10.1016/0009-2614(82)80012-2. 96. Fuentealba, P.; Stoll, H.; von Szentpaly, L.; Schwerdtfeger, P.; Preuss, H. On the Reliability of Semi-Empirical Pseudopotentials: Simulation of Hartree-Fock and Dirac-Fock Results. J. Phys. B Atomic Mol. Phys. 1983, 16 (11), L323–L328. https://doi.org/10.1088/0022-3700/16/11/001. 97. von Szentpály, L.; Fuentealba, P.; Preuss, H.; Stoll, H. Pseudopotential Calculations on Rb +2, Cs +2, RbH +, CsH + and the Mixed Alkali Dimer Ions. Chem. Phys. Lett. 1982, 93 (6), 555–559. https://doi.org/10.1016/0009-2614(82)83728-7. 98. Pantazis, D. A.; Neese, F. All-Electron Scalar Relativistic Basis Sets for the Lanthanides. J. Chem. Theory Comput. 2009, 5 (9), 2229–2238. https://doi.org/10.1021/ ct900090f. 99. Pantazis, D. A.; Neese, F. All-Electron Scalar Relativistic Basis Sets for the 6p Elements. Theor. Chem. Accounts 2012, 131 (11), 1292. https://doi.org/10.1007/s00214-0121292-x. 100. Suryanarayana, P. On Nearsightedness in Metallic Systems for O(N) Density Functional Theory Calculations: A Case Study on Aluminum. Chem. Phys. Lett. 2017, 679, 146–151. https://doi.org/10.1016/j.cplett.2017.04.095. 101. Xue, H.-T.; Boschetto, G.; Krompiec, M.; Morse, G. E.; Tang, F.-L.; Skylaris, C.-K. Linear-Scaling Density Functional Simulations of the Effect of Crystallographic Structure on the Electronic and Optical Properties of Fullerene Solvates. Phys. Chem. Chem. Phys. 2017, 19 (7), 5617–5628. https://doi.org/10.1039/C6CP08165G. 102. Freitag, L.; Knecht, S.; Angeli, C.; Reiher, M. Multireference Perturbation Theory with Cholesky Decomposition for the Density Matrix Renormalization Group. J. Chem. Theory Comput. 2017, 13 (2), 451–459. https://doi.org/10.1021/acs.jctc.6b00778. 103. Nakatani, N.; Guo, S. Density Matrix Renormalization Group (DMRG) Method as a Common Tool for Large Active-Space CASSCF/CASPT2 Calculations. J. Chem. Phys. 2017, 146 (9), 094102https://doi.org/10.1063/1.4976644. 104. Sayfutyarova, E. R.; Chan, G. K.-L. A State Interaction Spin-Orbit Coupling Density Matrix Renormalization Group Method. J. Chem. Phys. 2016, 144 (23), 234301. https://doi. org/10.1063/1.4953445. 105. Morgante, P.; Peverati, R. The Devil in the Details: What Everybody Should Know When Running DFT Calculations; https://doi.org/10.26434/chemrxiv.10187756.v1. 106. Boguslawski, K.; Jacob, C. R.; Reiher, M. Can DFT Accurately Predict Spin Densities? Analysis of Discrepancies in Iron Nitrosyl Complexes. J. Chem. Theory Comput. 2011, 7 (9), 2740–2752. https://doi.org/10.1021/ct1006218. 107. Radon, M.; Broclawik, E.; Pierloot, K. Electronic Structure of Selected {FeNO}7 Complexes in Heme and Non-Heme Architectures: A Density Functional and Multireference Ab Initio Study. J. Phys. Chem. B 2010, 114 (3), 1518–1528. https://doi.org/10.1021/jp910220r. 108. Roos, B. O.; Taylor, P. R.; Sigbahn, P. E. M. A Complete Active Space SCF Method (CASSCF) Using a Density Matrix Formulated Super-CI Approach. Chem. Phys. 1980, 48 (2), 157–173. https://doi.org/10.1016/0301-0104(80)80045-0. 109. Smith, J. E. T.; Mussard, B.; Holmes, A. A.; Sharma, S. Cheap and Near Exact CASSCF with Large Active Spaces. J. Chem. Theory Comput. 2017, 13 (11), 5468–5478. https:// doi.org/10.1021/acs.jctc.7b00900. 110. Malmqvist, P.A˚ .; Pierloot, K.; Shahi, A. R. M.; Cramer, C. J.; Gagliardi, L. The Restricted Active Space Followed by Second-Order Perturbation Theory Method: Theory and Application to the Study of CuO2 and Cu2O2 Systems. J. Chem. Phys. 2008, 128 (20), 204109. https://doi.org/10.1063/1.2920188. 111. Malmqvist, P. A.; Rendell, A.; Roos, B. O. The Restricted Active Space Self-Consistent-Field Method, Implemented With a Split Graph Unitary Group Approach. J. Phys. Chem. 1990, 94 (14), 5477–5482. https://doi.org/10.1021/j100377a011. 112. Ma, D.; Li Manni, G.; Gagliardi, L. The Generalized Active Space Concept in Multiconfigurational Self-Consistent Field Methods. J. Chem. Phys. 2011, 135 (4), 044128https:// doi.org/10.1063/1.3611401. 113. Neese, F.; Petrenko, T.; Ganyushin, D.; Olbrich, G. Advanced Aspects of Ab Initio Theoretical Optical Spectroscopy of Transition Metal Complexes: Multiplets, Spin-Orbit Coupling and Resonance Raman Intensities. Coord. Chem. Rev. 2007, 251 (3–4), 288–327. https://doi.org/10.1016/j.ccr.2006.05.019. 114. Ridley, J.; Zerner, M. An Intermediate Neglect of Differential Overlap Technique for Spectroscopy: Pyrrole and the Azines. Theor. Chim. Acta 1973, 32 (2), 111–134. https://doi. org/10.1007/BF00528484. 115. Zerner, M. C. Electronic Structure Theory for Transition Metal Systems: A Survey. In Metal-Ligand Interactions: From Atoms, to Clusters, to Surfaces; Salahub, D. R., Russo, N., Eds.; NATO ASI Series Springer Netherlands: Dordrecht, 1992;; pp 101–123. https://doi.org/10.1007/978-94-011-2822-3_5. 116. Zerner, M. C. Intermediate Neglect of Differential Overlap Calculations on the Electronic Spectra of Transition Metal Complexes. In Metal-Ligand Interactions: Structure and Reactivity; Russo, N., Salahub, D. R., Eds.; NATO ASI Series Springer Netherlands: Dordrecht, 1996;; pp 493–531. https://doi.org/10.1007/978-94-009-0155-1_18. 117. Zerner, M. C.; Loew, G. H.; Kirchner, R. F.; Mueller-Westerhoff, U. T. An Intermediate Neglect of Differential Overlap Technique for Spectroscopy of Transition-Metal Complexes. Ferrocene. J. Am. Chem. Soc. 1980, 102 (2), 589–599. https://doi.org/10.1021/ja00522a025. 118. Greene, S. N.; Richards, N. G. J. Electronic Structure, Bonding, Spectroscopy and Energetics of Fe-Dependent Nitrile Hydratase Active-Site Models. Inorg. Chem. 2006, 45 (1), 17–36. https://doi.org/10.1021/ic050965p. 119. Praneeth, V. K. K.; Näther, C.; Peters, G.; Lehnert, N. Spectroscopic Properties and Electronic Structure of Five- and Six-Coordinate Iron(II) Porphyrin NO Complexes: Effect of the Axial N-Donor Ligand. Inorg. Chem. 2006, 45 (7), 2795–2811. https://doi.org/10.1021/ic050865j. 120. Ottonelli, M.; Izzo, G. M. M.; Rizzo, F.; Musso, G.; Dellepiane, G.; Tubino, R. Semiempirical Study of the Electronic and Optical Properties of the Er(8-Hydroxyquinolinate)3 Complex. J. Phys. Chem. B 2005, 109 (41), 19249–19256. https://doi.org/10.1021/jp053314n. 121. O’Brien, T. A. Spin −Orbit Effects in the Ground States of Singly Positive and Neutral V2, VNb, and Nb2: INDO/S and Empirical Model Calculations. J. Phys. Chem. A 2004, 108 (23), 5016–5025. https://doi.org/10.1021/jp014108s. 122. Jackson, T. A.; Karapetian, A.; Miller, A.-F.; Brunold, T. C. Spectroscopic and Computational Studies of the Azide-Adduct of Manganese Superoxide Dismutase: Definitive Assignment of the Ligand Responsible for the Low-Temperature Thermochromism. J. Am. Chem. Soc. 2004, 126 (39), 12477–12491. https://doi.org/10.1021/ja0482583. 123. Greene, S. N.; Richards, N. G. J. Theoretical Investigations of the Electronic Structure and Spectroscopy of Mononuclear, Non-Heme {Fe −NO}6 Complexes. Inorg. Chem. 2004, 43 (22), 7030–7041. https://doi.org/10.1021/ic0499695. 124. Neese, F. A Spectroscopy Oriented Configuration Interaction Procedure. J. Chem. Phys. 2003, 119 (18), 9428–9443. https://doi.org/10.1063/1.1615956.

Computational Methods in Organometallic Chemistry

205

125. Neese, F. Theoretical Spectroscopy of Model-Nonheme [Fe(IV)OL5]2 + Complexes in Their Lowest Triplet and Quintet States Using Multireference Ab Initio and Density Functional Theory Methods. J. Inorg. Biochem. 2006, 100 (4), 716–726. https://doi.org/10.1016/j.jinorgbio.2006.01.020. 126. Wanko, M.; Hoffmann, M.; Strodel, P.; Koslowski, A.; Thiel, W.; Neese, F.; Frauenheim, T.; Elstner, M. Calculating Absorption Shifts for Retinal Proteins: Computational Challenges. J. Phys. Chem. B 2005, 109 (8), 3606–3615. https://doi.org/10.1021/jp0463060. 127. Schöneboom, J. C.; Neese, F.; Thiel, W. Toward Identification of the Compound I Reactive Intermediate in Cytochrome P450 Chemistry: A QM/MM Study of Its EPR and Mössbauer Parameters. J. Am. Chem. Soc. 2005, 127 (16), 5840–5853. https://doi.org/10.1021/ja0424732. 128. Blanchard, S.; Neese, F.; Bothe, E.; Bill, E.; Weyhermüller, T.; Wieghardt, K. Square Planar vs Tetrahedral Coordination in Diamagnetic Complexes of Nickel(II) Containing Two Bidentate p-Radical Monoanions. Inorg. Chem. 2005, 44 (10), 3636–3656. https://doi.org/10.1021/ic040117e. 129. Fouqueau, A.; Mer, S.; Casida, M. E.; Lawson Daku, L. M.; Hauser, A.; Mineva, T.; Neese, F. Comparison of Density Functionals for Energy and Structural Differences between the High- [5T2g: (T2g)4(Eg)2] and Low- [1A1g: (T2g)6(Eg)0] Spin States of the Hexaquoferrous Cation [Fe(H2O)6]2 +. J. Chem. Phys. 2004, 120 (20), 9473–9486. https://doi. org/10.1063/1.1710046. 130. Enemark, J. H.; Feltham, R. D. Principles of Structure, Bonding, and Reactivity for Metal Nitrosyl Complexes. Coord. Chem. Rev. 1974, 13 (4), 339–406. https://doi.org/ 10.1016/S0010-8545(00)80259-3. 131. Nardis, S.; Stefanelli, M.; Mohite, P.; Pomarico, G.; Tortora, L.; Manowong, M.; Chen, P.; Kadish, K. M.; Fronczek, F. R.; McCandless, G. T.; Smith, K. M.; Paolesse, R. b-Nitro Derivatives of Iron Corrolates. Inorg. Chem. 2012, 51 (6), 3910–3920. https://doi.org/10.1021/ic3002459. 132. Simkhovich, L.; Goldberg, I.; Gross, Z. Iron(III) and Iron(IV) Corroles: Synthesis, Spectroscopy, Structures, and No Indications for Corrole Radicals. Inorg. Chem. 2002, 41 (21), 5433–5439. https://doi.org/10.1021/ic020118b. 133. Simkhovich, L.; Mahammed, A.; Goldberg, I.; Gross, Z. Synthesis and Characterization of Germanium, Tin, Phosphorus, Iron, and Rhodium Complexes of Tris(Pentafluorophenyl) Corrole, and the Utilization of the Iron and Rhodium Corroles as Cyclopropanation Catalysts. Chem. Eur. J. 2001, 7 (5), 1041–1055. https://doi.org/10.1002/1521-3765 (20010302)7:53.0.CO;2-8. 134. Autret, M.; Will, S.; Caemelbecke, E. V.; Lex, J.; Gisselbrecht, J.-P.; Gross, M.; Vogel, E.; Kadish, K. M. Synthesis and Electrochemistry of Iron(III) Corroles Containing a Nitrosyl Axial Ligand. Spectral Characterization of [(OEC)FeIII(NO)]n Where n ¼ 0, 1, 2, or -1 and OEC Is the Trianion of 2,3,7,8,12,13,17,18-Octaethylcorrole. J. Am. Chem. Soc. 1994, 116 (20), 9141–9149. https://doi.org/10.1021/ja00099a032. 135. Ganguly, S.; Giles, L. J.; Thomas, K. E.; Sarangi, R.; Ghosh, A. Ligand Noninnocence in Iron Corroles: Insights from Optical and X-Ray Absorption Spectroscopies and Electrochemical Redox Potentials. Chem. Eur. J. 2017, 23 (60), 15098–15106. https://doi.org/10.1002/chem.201702621. 136. Norheim, H.-K.; Capar, J.; Einrem, R. F.; Gagnon, K. J.; Beavers, C. M.; Vazquez-Lima, H.; Ghosh, A. Ligand Noninnocence in FeNO Corroles: Insights From b-Octabromocorrole Complexes. Dalton Trans. 2015, 45 (2), 681–689. https://doi.org/10.1039/C5DT03947A. 137. Vazquez-Lima, H.; Norheim, H.-K.; Einrem, R. F.; Ghosh, A. Cryptic Noninnocence: FeNO Corroles in a New Light. Dalton Trans. 2015, 44 (22), 10146–10151. https://doi.org/ 10.1039/C5DT01495F. 138. Pierloot, K.; Phung, Q. M.; Ghosh, A. Electronic Structure of Neutral and Anionic Iron–Nitrosyl Corrole. A Multiconfigurational and Density Matrix Renormalization Group Investigation. Inorg. Chem. 2020, 59 (16), 11493–11502. https://doi.org/10.1021/acs.inorgchem.0c01312. 139. Rahman, M. H.; Ryan, M. D.; Vazquez-Lima, H.; Alemayehu, A.; Ghosh, A. Infrared Spectroelectrochemistry of Iron-Nitrosyl Triarylcorroles. Implications for Ligand Noninnocence. Inorg. Chem. 2020, 59 (5), 3232–3238. https://doi.org/10.1021/acs.inorgchem.9b03613. 140. de Groot, F. Multiplet Effects in X-Ray Spectroscopy. Coord. Chem. Rev. 2005, 249 (1), 31–63. https://doi.org/10.1016/j.ccr.2004.03.018. 141. Cowan, R. D. The Theory of Atomic Structure and Spectra; University of California Press, 1981. 142. Turner, S.; Lazar, S.; Freitag, B.; Egoavil, R.; Verbeeck, J.; Put, S.; Strauven, Y.; Tendeloo, G. V. High Resolution Mapping of Surface Reduction in Ceria Nanoparticles. Nanoscale 2011, 3 (8), 3385–3390. https://doi.org/10.1039/C1NR10510H. 143. Kucheyev, S. O.; Clapsaddle, B. J.; Wang, Y. M.; van Buuren, T.; Hamza, A. V. Electronic Structure of Nanoporous Ceria from X-Ray Absorption Spectroscopy and Atomic Multiplet Calculations. Phys. Rev. B 2007, 76 (23), 235420. https://doi.org/10.1103/PhysRevB.76.235420. 144. Stavitski, E.; de Groot, F. M. F. The CTM4XAS Program for EELS and XAS Spectral Shape Analysis of Transition Metal L Edges. Micron 2010, 41 (7), 687–694. https://doi.org/ 10.1016/j.micron.2010.06.005. 145. Löble, M. W.; Keith, J. M.; Altman, A. B.; Stieber, S. C. E.; Batista, E. R.; Boland, K. S.; Conradson, S. D.; Clark, D. L.; Lezama Pacheco, J.; Kozimor, S. A.; Martin, R. L.; Minasian, S. G.; Olson, A. C.; Scott, B. L.; Shuh, D. K.; Tyliszczak, T.; Wilkerson, M. P.; Zehnder, R. A. Covalency in Lanthanides. An X-Ray Absorption Spectroscopy and Density Functional Theory Study of LnCl 6 x – ( x ¼ 3, 2). J. Am. Chem. Soc. 2015, 137 (7), 2506–2523. https://doi.org/10.1021/ja510067v. 146. Minasian, S. G.; Batista, E. R.; Booth, C. H.; Clark, D. L.; Keith, J. M.; Kozimor, S. A.; Lukens, W. W.; Martin, R. L.; Shuh, D. K.; Stieber, S. C. E.; Tylisczcak, T.; Wen, X. Quantitative Evidence for Lanthanide-Oxygen Orbital Mixing in CeO2, PrO2, and TbO2. J. Am. Chem. Soc. 2017, 139 (49), 18052–18064. https://doi.org/10.1021/ jacs.7b10361. 147. Smiles, D. E.; Batista, E. R.; Booth, C. H.; Clark, D. L.; Keith, J. M.; Kozimor, S. A.; Martin, R. L.; Minasian, S. G.; Shuh, D. K.; Stieber, S. C. E.; Tyliszczak, T. The Duality of Electron Localization and Covalency in Lanthanide and Actinide Metallocenes. Chem. Sci. 2020, 11 (10), 2796–2809. https://doi.org/10.1039/C9SC06114B. 148. de Groot, F.; Kotani, A.; Kotani, A. Core Level Spectroscopy of Solids; CRC Press, 2008https://doi.org/10.1201/9781420008425. 149. Dreuw, A.; Weisman, J. L.; Head-Gordon, M. Long-Range Charge-Transfer Excited States in Time-Dependent Density Functional Theory Require Non-Local Exchange. J. Chem. Phys. 2003, 119 (6), 2943–2946. https://doi.org/10.1063/1.1590951. 150. Dreuw, A.; Head-Gordon, M. Failure of Time-Dependent Density Functional Theory for Long-Range Charge-Transfer Excited States: The Zincbacteriochlorin −Bacteriochlorin and Bacteriochlorophyll − Spheroidene Complexes. J. Am. Chem. Soc. 2004, 126 (12), 4007–4016. https://doi.org/10.1021/ja039556n. 151. Nemykin, V. N.; Hadt, R. G.; Belosludov, R. V.; Mizuseki, H.; Kawazoe, Y. Influence of Molecular Geometry, Exchange-Correlation Functional, and Solvent Effects in the Modeling of Vertical Excitation Energies in Phthalocyanines Using Time-Dependent Density Functional Theory (TDDFT) and Polarized Continuum Model TDDFT Methods: Can Modern Computational Chemistry Methods Explain Experimental Controversies?J. Phys. Chem. A 2007, 111 (50), 12901–12913. https://doi.org/10.1021/jp0759731. 152. Gritsenko, O.; Baerends, E. J. Asymptotic Correction of the Exchange–Correlation Kernel of Time-Dependent Density Functional Theory for Long-Range Charge-Transfer Excitations. J. Chem. Phys. 2004, 121 (2), 655–660. https://doi.org/10.1063/1.1759320. 153. Schipper, P. R. T.; Gritsenko, O. V.; van Gisbergen, S. J. A.; Baerends, E. J. Molecular Calculations of Excitation Energies and (Hyper)Polarizabilities with a Statistical Average of Orbital Model Exchange-Correlation Potentials. J. Chem. Phys. 2000, 112 (3), 1344–1352. https://doi.org/10.1063/1.480688. 154. Neugebauer, J.; Gritsenko, O.; Baerends, E. J. Assessment of a Simple Correction for the Long-Range Charge-Transfer Problem in Time-Dependent Density-Functional Theory. J. Chem. Phys. 2006, 124 (21), 214102. https://doi.org/10.1063/1.2197829. 155. Peach, M. J. G.; Benfield, P.; Helgaker, T.; Tozer, D. J. Excitation Energies in Density Functional Theory: An Evaluation and a Diagnostic Test. J. Chem. Phys. 2008, 128 (4), 044118https://doi.org/10.1063/1.2831900. 156. Ziegler, T.; Seth, M.; Krykunov, M.; Autschbach, J.; Wang, F. Is Charge Transfer Transitions Really Too Difficult for Standard Density Functionals or Are They Just a Problem for Time-Dependent Density Functional Theory Based on a Linear Response Approach. J. Mol. Struct. THEOCHEM 2009, 914 (1), 106–109. https://doi.org/10.1016/j. theochem.2009.04.021. 157. Cavillot, V.; Champagne, B. Time-Dependent Density Functional Theory Simulation of UV/Visible Absorption Spectra of Zirconocene Catalysts. Chem. Phys. Lett. 2002, 354 (5), 449–457. https://doi.org/10.1016/S0009-2614(02)00161-6. 158. van Gisbergen, S. J. A.; Rosa, A.; Ricciardi, G.; Baerends, E. J. Time-Dependent Density Functional Calculations on the Electronic Absorption Spectrum of Free Base Porphin. J. Chem. Phys. 1999, 111 (6), 2499–2506. https://doi.org/10.1063/1.479617.

206

Computational Methods in Organometallic Chemistry

159. Sundholm, D. Density Functional Theory Calculations of the Visible Spectrum of Chlorophyll A. Chem. Phys. Lett. 1999, 302 (5), 480–484. https://doi.org/10.1016/S00092614(99)00194-3. 160. Adamo, C.; Barone, V. Inexpensive and Accurate Predictions of Optical Excitations in Transition-Metal Complexes: The TDDFT/PBE0 Route. Theor. Chem. Accounts 2000, 105 (2), 169–172. https://doi.org/10.1007/s002140000202. 161. Full, J.; González, L.; Daniel, C. A CASSCF/CASPT2 and TD-DFT Study of the Low-Lying Excited States of Η5-CpMn(CO)3. J. Phys. Chem. A 2001, 105 (1), 184–189. https:// doi.org/10.1021/jp002042f. 162. Boulet, P.; Chermette, H.; Daul, C.; Gilardoni, F.; Rogemond, F.; Weber, J.; Zuber, G. Absorption Spectra of Several Metal Complexes Revisited by the Time-Dependent Density-Functional Theory-Response Theory Formalism. J. Phys. Chem. A 2001, 105 (5), 885–894. https://doi.org/10.1021/jp003041q. 163. Seidu, I.; Krykunov, M.; Ziegler, T. Applications of Time-Dependent and Time-Independent Density Functional Theory to Electronic Transitions in Tetrahedral D0 Metal Oxides. J. Chem. Theory Comput. 2015, 11 (9), 4041–4053. https://doi.org/10.1021/acs.jctc.5b00298. 164. Freitag, L.; González, L. Theoretical Spectroscopy and Photodynamics of a Ruthenium Nitrosyl Complex. Inorg. Chem. 2014, 53 (13), 6415–6426. https://doi.org/10.1021/ ic500283y. 165. Becke, A. D. Density-functional Thermochemistry. III. The Role of Exact Exchange. J. Chem. Phys. 1993, 98 (7), 5648–5652. https://doi.org/10.1063/1.464913. 166. Andrae, D.; Häußermann, U.; Dolg, M.; Stoll, H.; Preuß, H. Energy-Adjustedab Initio Pseudopotentials for the Second and Third Row Transition Elements. Theor. Chim. Acta 1990, 77 (2), 123–141. https://doi.org/10.1007/BF01114537. 167. Eichkorn, K.; Treutler, O.; Öhm, H.; Häser, M.; Ahlrichs, R. Auxiliary Basis Sets to Approximate Coulomb Potentials. Chem. Phys. Lett. 1995, 240 (4), 283–290. https://doi.org/ 10.1016/0009-2614(95)00621-A. 168. Vahtras, O.; Almlöf, J.; Feyereisen, M. W. Integral Approximations for LCAO-SCF Calculations. Chem. Phys. Lett. 1993, 213 (5), 514–518. https://doi.org/10.1016/0009-2614 (93)89151-7. 169. Finley, J.; Malmqvist, P.-A˚ .; Roos, B. O.; Serrano-Andrés, L. The Multi-State CASPT2 Method. Chem. Phys. Lett. 1998, 288 (2), 299–306. https://doi.org/10.1016/S00092614(98)00252-8. 170. Bytheway, I.; Wong, M. W. The Prediction of Vibrational Frequencies of Inorganic Molecules Using Density Functional Theory. Chem. Phys. Lett. 1998, 282 (3), 219–226. https://doi.org/10.1016/S0009-2614(97)01281-5. 171. Andrews, L.; Citra, A. Infrared Spectra and Density Functional Theory Calculations on Transition Metal Nitrosyls. Vibrational Frequencies of Unsaturated Transition Metal Nitrosyls. Chem. Rev. 2002, 102 (4), 885–912. https://doi.org/10.1021/cr0000729. 172. Rauhut, G.; Jarzecki, A. A.; Pulay, P. Density Functional Based Vibrational Study of Conformational Isomers: Molecular Rearrangement of Benzofuroxan. J. Comput. Chem. 1997, 18 (4), 489–500. https://doi.org/10.1002/(SICI)1096-987X(199703)18:43.0.CO;2-P. 173. Lanucara, F.; Chiavarino, B.; Crestoni, M. E.; Scuderi, D.; Sinha, R. K.; Maıtre, P.; Fornarini, S. Naked Five-Coordinate FeIII(NO) Porphyrin Complexes: Vibrational and Reactivity Features. Inorg. Chem. 2011, 50 (10), 4445–4452. https://doi.org/10.1021/ic200073v. 174. Rassolov, V. A.; Pople, J. A.; Ratner, M. A.; Windus, T. L. 6-31G Basis Set for Atoms K Through Zn. J. Chem. Phys. 1998, 109 (4), 1223–1229. https://doi.org/ 10.1063/1.476673. 175. MacAleese, L.; Maître, P. Infrared Spectroscopy of Organometallic Ions in the Gas Phase: From Model to Real World Complexes. Mass Spectrom. Rev. 2007, 26 (4), 583–605. https://doi.org/10.1002/mas.20138. 176. Oomens, J.; van Roij, A. J. A.; Meijer, G.; von Helden, G. Gas-Phase Infrared Photodissociation Spectroscopy of Cationic Polyaromatic Hydrocarbons. ApJs 2000, 542 (1), 404. https://doi.org/10.1086/309545. 177. Polfer, N. C.; Oomens, J. Reaction Products in Mass Spectrometry Elucidated With Infrared Spectroscopy. Phys. Chem. Chem. Phys. 2007, 9 (29), 3804–3817. https://doi.org/ 10.1039/B702993B. 178. Chiavarino, B.; Crestoni, M. E.; Fornarini, S.; Lemaire, J.; Maître, P.; MacAleese, L. p-Complex Structure of Gaseous Benzene − NO Cations Assayed by IR Multiple Photon Dissociation Spectroscopy. J. Am. Chem. Soc. 2006, 128 (38), 12553–12561. https://doi.org/10.1021/ja0637548. 179. Lodewyk, M. W.; Siebert, M. R.; Tantillo, D. J. Computational Prediction of 1 H and 13 C Chemical Shifts: A Useful Tool for Natural Product, Mechanistic, and Synthetic Organic Chemistry. Chem. Rev. 2012, 112 (3), 1839–1862. https://doi.org/10.1021/cr200106v. 180. Machácek, J.; Bühl, M.; Fanfrlík, J.; Hnyk, D. Nuclear Magnetic Shielding of Monoboranes: Calculation and Assessment of 11B NMR Chemical Shifts in Planar BX3 and in Tetrahedral [BX4]− Systems. J. Phys. Chem. A 2017, 121 (50), 9631–9637. https://doi.org/10.1021/acs.jpca.7b09831. 181. Contreras, R. H.; Llorente, T.; Pagola, G. I.; Bustamante, M. G.; Pasqualini, E. E.; Melo, J. I.; Tormena, C. F. Qualitative Study of Substituent Effects on NMR 15N and 17O Chemical Shifts. J. Phys. Chem. A 2009, 113 (36), 9874–9880. https://doi.org/10.1021/jp901926p. 182. Zhu, J.; Kurahashi, T.; Fujii, H.; Wu, G. Solid-State 17O NMR and Computational Studies of Terminal Transition Metal Oxo Compounds. Chem. Sci. 2012, 3 (2), 391–397. https://doi.org/10.1039/C1SC00725D. 183. Moroz, I. B.; Larmier, K.; Liao, W.-C.; Copéret, C. Discerning g-Alumina Surface Sites With Nitrogen-15 Dynamic Nuclear Polarization Surface Enhanced NMR Spectroscopy of Adsorbed Pyridine. J. Phys. Chem. C 2018, 122 (20), 10871–10882. https://doi.org/10.1021/acs.jpcc.8b01823. 184. Ehinger, C.; Gordon, C. P.; Copéret, C. Oxygen Transfer in Electrophilic Epoxidation Probed by 17O NMR: Differentiating Between Oxidants and Role of Spectator Metal Oxo. Chem. Sci. 2019, 10 (6), 1786–1795. https://doi.org/10.1039/C8SC04868A. 185. Lam, E.; Copéret, C. Understanding Trends in 27Al Chemical Shifts and Quadrupolar Coupling Constants in Chloroalkyl Aluminum [AlClx(Me)3 − x]n ¼ 1 or 2 Compounds. Helv. Chim. Acta 2018, 101 (9), e1800120https://doi.org/10.1002/hlca.201800120. 186. Moroz, I. B.; Florian, P.; Viger-Gravel, J.; Gordon, C. P.; Lesage, A.; Copéret, C. Silica-Grafted Tris(Neopentyl)Aluminum: A Monomeric Aluminum Solid Co-Catalyst for Efficient Nickel-Catalyzed Ethene Dimerization. Angew. Chem. Int. Ed. 2020, 59 (37), 16167–16172. https://doi.org/10.1002/anie.202006285. 187. Cavalieri, J. D.; West, R.; Duchamp, J. C.; Zilm, K. W. Unusual Silicon-29 Chemical-Shift Anisotropies in Three-Membered Rings. J. Am. Chem. Soc. 1993, 115 (9), 3770–3771. https://doi.org/10.1021/ja00062a051. 188. West, R.; Cavalieri, J. D.; Buffy, J. J.; Fry, C.; Zilm, K. W.; Duchamp, J. C.; Kira, M.; Iwamoto, T.; Müller, T.; Apeloig, Y. A Solid-State 1NMR and Theoretical Study of the Chemical Bonding in Disilenes. J. Am. Chem. Soc. 1997, 119 (21), 4972–4976. https://doi.org/10.1021/ja963921b. 189. Buffy, J. J.; West, R.; Bendikov, M.; Apeloig, Y. Chemical Shielding Tensors for a Silicon −Carbon Double Bond. J. Am. Chem. Soc. 2001, 123 (5), 978–979. https://doi.org/ 10.1021/ja003389z. 190. Kravchenko, V.; Kinjo, R.; Sekiguchi, A.; Ichinohe, M.; West, R.; Balazs, Y. S.; Schmidt, A.; Karni, M.; Apeloig, Y. Solid-State 29Si NMR Study of RSiSiR: A Tool for Analyzing the Nature of the Si −Si Bond. J. Am. Chem. Soc. 2006, 128 (45), 14472–14473. https://doi.org/10.1021/ja065817s. 191. Duchamp, J. C.; Pakulski, M.; Cowley, A. H.; Zilm, K. W. Nature of the Carbon-Phosphorus Double Bond and the Carbon-Phosphorus Triple Bond as Studied by Solid-State NMR. J. Am. Chem. Soc. 1990, 112 (19), 6803–6809. https://doi.org/10.1021/ja00175a010. 192. Wu, G.; Rovnyak, D.; Johnson, M. J. A.; Zanetti, N. C.; Musaev, D. G.; Morokuma, K.; Schrock, R. R.; Griffin, R. G.; Cummins, C. C. Unusual 31P Chemical Shielding Tensors in Terminal Phosphido Complexes Containing a Phosphorus −Metal Triple Bond. J. Am. Chem. Soc. 1996, 118 (43), 10654–10655. https://doi.org/10.1021/ja960639w. 193. Eichele, K.; Wasylishen, R. E.; Corrigan, J. F.; Taylor, N. J.; Carty, A. J.; Feindel, K. W.; Bernard, G. M. Phosphorus Chemical Shift Tensors of Phosphido Ligands in Ruthenium Carbonyl Compounds: 31P NMR Spectroscopy of Single-Crystal and Powder Samples and Ab Initio Calculations. J. Am. Chem. Soc. 2002, 124 (7), 1541–1552. https://doi.org/ 10.1021/ja0122041. 194. Engl, P. S.; Santiago, C. B.; Gordon, C. P.; Liao, W.-C.; Fedorov, A.; Copéret, C.; Sigman, M. S.; Togni, A. Exploiting and Understanding the Selectivity of Ru-N-Heterocyclic Carbene Metathesis Catalysts for the Ethenolysis of Cyclic Olefins to a,o-Dienes. J. Am. Chem. Soc. 2017, 139 (37), 13117–13125. https://doi.org/10.1021/jacs.7b06947.

Computational Methods in Organometallic Chemistry

207

195. Vummaleti, S. V. C.; Nelson, D. J.; Poater, A.; Gómez-Suárez, A.; Cordes, D. B.; Slawin, A. M. Z.; Nolan, S. P.; Cavallo, L. What Can NMR Spectroscopy of Selenoureas and Phosphinidenes Teach Us About the p-Accepting Abilities of N-Heterocyclic Carbenes?Chem. Sci. 2015, 6 (3), 1895–1904. https://doi.org/10.1039/C4SC03264K. 196. Autschbach, J.; Zheng, S. Analyzing Pt Chemical Shifts Calculated From Relativistic Density Functional Theory Using Localized Orbitals: The Role of Pt 5d Lone Pairs. Magn. Reson. Chem. 2008, 46 (S1), S45–S55. https://doi.org/10.1002/mrc.2289. 197. Lätsch, L.; Lam, E.; Copéret, C. Electronegativity and Location of Anionic Ligands Drive Yttrium NMR for Molecular, Surface and Solid-State Structures. Chem. Sci. 2020, 11 (26), 6724–6735. https://doi.org/10.1039/D0SC02321C. 198. Gordon, C. P.; Raynaud, C.; Andersen, R. A.; Copéret, C.; Eisenstein, O. Carbon-13 NMR Chemical Shift: A Descriptor for Electronic Structure and Reactivity of Organometallic Compounds. Acc. Chem. Res. 2019, 52 (8), 2278–2289. https://doi.org/10.1021/acs.accounts.9b00225. 199. Gordon, C. P.; Yamamoto, K.; Liao, W.-C.; Allouche, F.; Andersen, R. A.; Copéret, C.; Raynaud, C.; Eisenstein, O. Metathesis Activity Encoded in the Metallacyclobutane Carbon-13 NMR Chemical Shift Tensors. ACS Cent. Sci. 2017, 3 (7), 759–768. https://doi.org/10.1021/acscentsci.7b00174. 200. Dolg, M.; Wedig, U.; Stoll, H.; Preuss, H. Energy-adjusted Ab Initio Pseudopotentials for the First Row Transition Elements. J. Chem. Phys. 1987, 86 (2), 866–872. https://doi. org/10.1063/1.452288. 201. Jensen, F. Unifying General and Segmented Contracted Basis Sets. Segmented Polarization Consistent Basis Sets. J. Chem. Theory Comput. 2014, 10 (3), 1074–1085. https:// doi.org/10.1021/ct401026a. 202. van Lenthe, E.; Baerends, E. J.; Snijders, J. G. Relativistic Regular Two-Component Hamiltonians. J. Chem. Phys. 1993, 99 (6), 4597–4610. https://doi.org/ 10.1063/1.466059. 203. van Lenthe, E.; Baerends, E. J.; Snijders, J. G. Relativistic Total Energy Using Regular Approximations. J. Chem. Phys. 1994, 101 (11), 9783–9792. https://doi.org/ 10.1063/1.467943. 204. van Lenthe, E.; Ehlers, A.; Baerends, E.-J. Geometry Optimizations in the Zero Order Regular Approximation for Relativistic Effects. J. Chem. Phys. 1999, 110 (18), 8943–8953. https://doi.org/10.1063/1.478813. 205. Rieger, A. L.; Rieger, P. H. Chemical Insights From EPR Spectra of Organometallic Radicals and Radical Ions. Organometallics 2004, 23 (2), 154–162. https://doi.org/10.1021/ om030565e. 206. Ames, W. M.; Larsen, S. C. DFT Calculations of EPR Parameters for Copper(II)-Exchanged Zeolites Using Cluster Models. J. Phys. Chem. A 2010, 114 (1), 589–594. https://doi. org/10.1021/jp907878h. 207. Hadt, R. G.; Nemykin, V. N.; Olsen, J. G.; Basu, P. Comparative Calculation of EPR Spectral Parameters in [MoVOX4]−, [MoVOX5]2 −, and [MoVOX4(H2O)]− Complexes. Phys. Chem. Chem. Phys. 2009, 11 (44), 10377–10384. https://doi.org/10.1039/B905554A. 208. Neese, F. Importance of Direct Spin − Spin Coupling and Spin-Flip Excitations for the Zero-Field Splittings of Transition Metal Complexes: A Case Study. J. Am. Chem. Soc. 2006, 128 (31), 10213–10222. https://doi.org/10.1021/ja061798a. 209. Rao, G.; Altman, A. B.; Brown, A. C.; Tao, L.; Stich, T. A.; Arnold, J.; Britt, R. D. Metal Bonding With 3d and 6d Orbitals: An EPR and ENDOR Spectroscopic Investigation of Ti 3+ –Al and Th 3+ –Al Heterobimetallic Complexes. Inorg. Chem. 2019, 58 (12), 7978–7988. https://doi.org/10.1021/acs.inorgchem.9b00720. 210. Staroverov, V. N.; Scuseria, G. E.; Tao, J.; Perdew, J. P. Erratum: “Comparative Assessment of a New Nonempirical Density Functional: Molecules and Hydrogen-Bonded Complexes” [J. Chem. Phys. 119, 12129 (2003)]. J. Chem. Phys. 2004, 121 (22), 11507. https://doi.org/10.1063/1.1795692. 211. Neese, F.; Wennmohs, F.; Hansen, A.; Becker, U. Efficient, Approximate and Parallel Hartree–Fock and Hybrid DFT Calculations. A ‘Chain-of-Spheres’ Algorithm for the Hartree–Fock Exchange. Chem. Phys. 2009, 356 (1), 98–109. https://doi.org/10.1016/j.chemphys.2008.10.036. 212. Altman, A. B.; Brown, A. C.; Rao, G.; Lohrey, T. D.; Britt, R. D.; Maron, L.; Minasian, S. G.; Shuh, D. K.; Arnold, J. Chemical Structure and Bonding in a Thorium(III)–Aluminum Heterobimetallic Complex. Chem. Sci. 2018, 9 (18), 4317–4324. https://doi.org/10.1039/C8SC01260A. 213. Bihlmayer, G. Density Functional Theory for Magnetism and Magnetic Anisotropy. In Handbook of Materials Modeling: Methods: Theory and Modeling; Andreoni, W., Yip, S., Eds.; Springer International Publishing: Cham, 2018;; pp 1–23. https://doi.org/10.1007/978-3-319-42913-7_73-1. 214. Kahn, O. Molecular Magnetism, 1st ed.; Wiley-VCH: New York, 1993. 215. Daul, C. Density Functional Theory Applied to the Excited States of Coordination Compounds. Int. J. Quantum Chem. 1994, 52 (4), 867–877. https://doi.org/10.1002/ qua.560520414. 216. Ovchinnikov, A. A.; Labanowski, J. K. Simple Spin Correction of Unrestricted Density-Functional Calculation. Phys. Rev. A 1996, 53 (6), 3946–3952. https://doi.org/10.1103/ PhysRevA.53.3946. 217. Gütlich, P.; Bill, E.; Trautwein, A. X. Mössbauer Spectroscopy and Transition Metal Chemistry: Fundamentals and Applications; Springer-Verlag: Berlin, Heidelberg, 2011https:// doi.org/10.1007/978-3-540-88428-6. 218. Woodruff, D. N.; Winpenny, R. E. P.; Layfield, R. A. Lanthanide Single-Molecule Magnets. Chem. Rev. 2013, 113 (7), 5110–5148. https://doi.org/10.1021/cr400018q. 219. Bogani, L.; Wernsdorfer, W. Molecular Spintronics Using Single-Molecule Magnets. Nat. Mater. 2008, 7 (3), 179–186. https://doi.org/10.1038/nmat2133. 220. Madhu, N. T.; Tang, J.-K.; Hewitt, I. J.; Clérac, R.; Wernsdorfer, W.; van Slageren, J.; Anson, C. E.; Powell, A. K. What Makes a Single Molecule Magnet?Polyhedron 2005, 24 (16), 2864–2869. https://doi.org/10.1016/j.poly.2005.03.015. 221. Shao, D.; Wang, X.-Y. Development of Single-Molecule Magnets{. Chin. J. Chem. 2020, 38 (9), 1005–1018. https://doi.org/10.1002/cjoc.202000090. 222. Graham, M. J.; Zadrozny, J. M.; Fataftah, M. S.; Freedman, D. E. Forging Solid-State Qubit Design Principles in a Molecular Furnace. Chem. Mater. 2017, 29 (5), 1885–1897. https://doi.org/10.1021/acs.chemmater.6b05433. 223. Liddle, S. T.; van Slageren, J. Improving F-Element Single Molecule Magnets. Chem. Soc. Rev. 2015, 44 (19), 6655–6669. https://doi.org/10.1039/C5CS00222B. 224. Pedersen, K. S.; Dreiser, J.; Weihe, H.; Sibille, R.; Johannesen, H. V.; Sørensen, M. A.; Nielsen, B. E.; Sigrist, M.; Mutka, H.; Rols, S.; Bendix, J.; Piligkos, S. Design of Single-Molecule Magnets: Insufficiency of the Anisotropy Barrier as the Sole Criterion. Inorg. Chem. 2015, 54 (15), 7600–7606. https://doi.org/10.1021/acs. inorgchem.5b01209. 225. Singh, S. K.; Cramer, C. J.; Gagliardi, L. Correlating Electronic Structure and Magnetic Anisotropy in Actinide Complexes [An(COT)2], AnIII/IV ¼ U, Np, and Pu. Inorg. Chem. 2020, 59 (10), 6815–6825. https://doi.org/10.1021/acs.inorgchem.0c00105. 226. Gaggioli, C. A.; Gagliardi, L. Theoretical Investigation of Plutonium-Based Single-Molecule Magnets. Inorg. Chem. 2018, 57 (14), 8098–8105. https://doi.org/10.1021/acs. inorgchem.8b00170. 227. Magnani, N.; Colineau, E.; Griveau, J.-C.; Apostolidis, C.; Walter, O.; Caciuffo, R. A Plutonium-Based Single-Molecule Magnet. Chem. Commun. 2014, 50 (60), 8171–8173. https://doi.org/10.1039/C4CC03400G. 228. Mössbauer, R. L. The Discovery of the Mössbauer Effect. Hyperfine Interact. 2000, 126 (1), 1–12. https://doi.org/10.1023/A:1012620106837. 229. Mössbauer, R. L. Kernresonanzfluoreszenz von Gammastrahlung in Ir191. Z. Phys. 1958, 151 (2), 124–143. https://doi.org/10.1007/BF01344210. 230. Mössbauer, R. L. Kernresonanzabsorption von Gammastrahlung in Ir191. Naturwissenschaften 1958, 45 (22), 538–539. https://doi.org/10.1007/BF00632050. 231. Gonser, U. From a Strange Effect to Mössbauer Spectroscopy. In Mössbauer Spectroscopy; Gonser, U., Ed.; Topics in Applied Physics Springer: Berlin, Heidelberg, 1975;; pp 1–51. https://doi.org/10.1007/3540071202_13. 232. Römelt, M.; Ye, S.; Neese, F. Calibration of Modern Density Functional Theory Methods for the Prediction of 57 Fe Mössbauer Isomer Shifts: Meta-GGA and Double-Hybrid Functionals. Inorg. Chem. 2009, 48 (3), 784–785. https://doi.org/10.1021/ic801535v. 233. Sandala, G. M.; Hopmann, K. H.; Ghosh, A.; Noodleman, L. Calibration of DFT Functionals for the Prediction of 57Fe Mössbauer Spectral Parameters in Iron–Nitrosyl and Iron–Sulfur Complexes: Accurate Geometries Prove Essential. J. Chem. Theory Comput. 2011, 7 (10), 3232–3247. https://doi.org/10.1021/ct200187d. 234. Bjornsson, R.; Neese, F.; DeBeer, S. Revisiting the Mössbauer Isomer Shifts of the FeMoco Cluster of Nitrogenase and the Cofactor Charge. Inorg. Chem. 2017, 56 (3), 1470–1477. https://doi.org/10.1021/acs.inorgchem.6b02540.

208

Computational Methods in Organometallic Chemistry

235. Gallenkamp, C.; Kramm, U. I.; Proppe, J.; Krewald, V. Calibration of Computational Mössbauer Spectroscopy to Unravel Active Sites in FeNC Catalysts for the Oxygen Reduction Reaction. Int. J. Quantum Chem. 2021, 121 (3), e26394https://doi.org/10.1002/qua.26394. 236. Bochevarov, A. D.; Friesner, R. A.; Lippard, S. J. Prediction of 57Fe Mössbauer Parameters by Density Functional Theory: A Benchmark Study. J. Chem. Theory Comput. 2010, 6 (12), 3735–3749. https://doi.org/10.1021/ct100398m. 237. Yano, J.; Yachandra, V. K. X-Ray Absorption Spectroscopy. Photosynth. Res. 2009, 102 (2–3), 241–254. https://doi.org/10.1007/s11120-009-9473-8. 238. Westre, T. E.; Kennepohl, P.; DeWitt, J. G.; Hedman, B.; Hodgson, K. O.; Solomon, E. I. A Multiplet Analysis of Fe K-Edge 1s ! 3d Pre-Edge Features of Iron Complexes. J. Am. Chem. Soc. 1997, 119 (27), 6297–6314. https://doi.org/10.1021/ja964352a. 239. Sarangi, R. X-Ray Absorption near-Edge Spectroscopy in Bioinorganic Chemistry: Application to M–O2 Systems. Coord. Chem. Rev. 2013, 257 (2), 459–472. https://doi.org/ 10.1016/j.ccr.2012.06.024. 240. DeBeer George, S.; Neese, F. Calibration of Scalar Relativistic Density Functional Theory for the Calculation of Sulfur K-Edge X-Ray Absorption Spectra. Inorg. Chem. 2010, 49 (4), 1849–1853. https://doi.org/10.1021/ic902202s. 241. Fronzoni, G.; Stener, M.; Reduce, A.; Decleva, P. Time-Dependent Density Functional Theory Calculations of Ligand K Edge and Metal L Edge X-Ray Absorption of a Series of Oxomolybdenum Complexes. J. Phys. Chem. A 2004, 108 (40), 8467–8477. https://doi.org/10.1021/jp047953u. 242. Casarin, M.; Finetti, P.; Vittadini, A.; Wang, F.; Ziegler, T. Spin −Orbit Relativistic Time-Dependent Density Functional Calculations of the Metal and Ligand Pre-Edge XAS Intensities of Organotitanium Complexes: TiCl4, Ti(Η5-C5H5)Cl3, and Ti(Η5-C5H5)2Cl2. J. Phys. Chem. A 2007, 111 (24), 5270–5279. https://doi.org/10.1021/jp071561g. 243. Minasian, S. G.; Keith, J. M.; Batista, E. R.; Boland, K. S.; Kozimor, S. A.; Martin, R. L.; Shuh, D. K.; Tyliszczak, T.; Vernon, L. J. Carbon K-Edge X-Ray Absorption Spectroscopy and Time-Dependent Density Functional Theory Examination of Metal–Carbon Bonding in Metallocene Dichlorides. J. Am. Chem. Soc. 2013, 135 (39), 14731–14740. https:// doi.org/10.1021/ja405844j. 244. Lee, K.; Wei, H.; Blake, A. V.; Donahue, C. M.; Keith, J. M.; Daly, S. R. Ligand K-Edge XAS, DFT, and TDDFT Analysis of Pincer Linker Variations in Rh( I ) PNP Complexes: Reactivity Insights From Electronic Structure. Dalton Trans. 2016, 45 (24), 9774–9785. https://doi.org/10.1039/C6DT00200E. 245. Sarangi, R.; DeBeer George, S.; Rudd, D. J.; Szilagyi, R. K.; Ribas, X.; Rovira, C.; Almeida, M.; Hodgson, K. O.; Hedman, B.; Solomon, E. I. Sulfur K-Edge X-Ray Absorption Spectroscopy as a Probe of Ligand − Metal Bond Covalency: Metal vs Ligand Oxidation in Copper and Nickel Dithiolene Complexes. J. Am. Chem. Soc. 2007, 129 (8), 2316–2326. https://doi.org/10.1021/ja0665949. 246. Baker, M. L.; Mara, M. W.; Yan, J. J.; Hodgson, K. O.; Hedman, B.; Solomon, E. I. K- and L-Edge X-Ray Absorption Spectroscopy (XAS) and Resonant Inelastic X-Ray Scattering (RIXS) Determination of Differential Orbital Covalency (DOC) of Transition Metal Sites. Coord. Chem. Rev. 2017, 345, 182–208. https://doi.org/10.1016/j.ccr.2017.02.004. 247. Spencer, L. P.; Yang, P.; Minasian, S. G.; Jilek, R. E.; Batista, E. R.; Boland, K. S.; Boncella, J. M.; Conradson, S. D.; Clark, D. L.; Hayton, T. W.; Kozimor, S. A.; Martin, R. L.; MacInnes, M. M.; Olson, A. C.; Scott, B. L.; Shuh, D. K.; Wilkerson, M. P. Tetrahalide Complexes of the [U(NR)2]2 + Ion: Synthesis, Theory, and Chlorine K-Edge X-Ray Absorption Spectroscopy. J. Am. Chem. Soc. 2013, 135 (6), 2279–2290. https://doi.org/10.1021/ja310575j. 248. Minasian, S. G.; Keith, J. M.; Batista, E. R.; Boland, K. S.; Clark, D. L.; Kozimor, S. A.; Martin, R. L.; Shuh, D. K.; Tyliszczak, T. New Evidence for 5f Covalency in Actinocenes Determined from Carbon K-Edge XAS and Electronic Structure Theory. Chem. Sci. 2013, 5 (1), 351–359. https://doi.org/10.1039/C3SC52030G. 249. DeBeer George, S.; Petrenko, T.; Neese, F. Prediction of Iron K-Edge Absorption Spectra Using Time-Dependent Density Functional Theory {. J. Phys. Chem. A 2008, 112 (50), 12936–12943. https://doi.org/10.1021/jp803174m. 250. DeBeer George, S.; Petrenko, T.; Neese, F. Time-Dependent Density Functional Calculations of Ligand K-Edge X-Ray Absorption Spectra. Inorg. Chim. Acta 2008, 361 (4), 965–972. https://doi.org/10.1016/j.ica.2007.05.046. 251. Ray, K.; DeBeer George, S.; Solomon, E. I.; Wieghardt, K.; Neese, F. Description of the Ground-State Covalencies of the Bis(Dithiolato) Transition-Metal Complexes From X-Ray Absorption Spectroscopy and Time-Dependent Density-Functional Calculations. Chem. Eur. J. 2007, 13 (10), 2783–2797. https://doi.org/10.1002/chem.200601425. 252. Roe, A. L.; Schneider, D. J.; Mayer, R. J.; Pyrz, J. W.; Widom, J.; Que, L. X-Ray Absorption Spectroscopy of Iron-Tyrosinate Proteins. J. Am. Chem. Soc. 1984, 106 (6), 1676–1681. https://doi.org/10.1021/ja00318a021. 253. Arrio, M.-A.; Rossano, S.; Brouder, C.; Galoisy, L.; Calas, G. Calculation of Multipole Transitions at the Fe K Pre-Edge through p-d Hybridization in the Ligand Field Multiplet Model. EPL 2000, 51 (4), 454. https://doi.org/10.1209/epl/i2000-00515-8. 254. Chandrasekaran, P.; Stieber, S. C. E.; Collins, T. J.; Lawrence Que, J.; Neese, F.; DeBeer, S. Prediction of High-Valent Iron K-Edge Absorption Spectra by Time-Dependent Density Functional Theory. Dalton Trans. 2011, 40 (42), 11070–11079. https://doi.org/10.1039/C1DT11331C. 255. Becke, A. D. Density-Functional Exchange-Energy Approximation With Correct Asymptotic Behavior. Phys. Rev. A 1988, 38 (6), 3098–3100. https://doi.org/10.1103/ PhysRevA.38.3098. 256. Neese, F. Prediction and Interpretation of the 57Fe Isomer Shift in Mössbauer Spectra by Density Functional Theory. Inorg. Chim. Acta 2002, 337, 181–192. https://doi.org/ 10.1016/S0020-1693(02)01031-9. 257. Maganas, D.; DeBeer, S.; Neese, F. Pair Natural Orbital Restricted Open-Shell Configuration Interaction (PNO-ROCIS) Approach for Calculating X-Ray Absorption Spectra of Large Chemical Systems. J. Phys. Chem. A 2018, 122 (5), 1215–1227. https://doi.org/10.1021/acs.jpca.7b10880. 258. Minasian, S. G.; Keith, J. M.; Batista, E. R.; Boland, K. S.; Clark, D. L.; Conradson, S. D.; Kozimor, S. A.; Martin, R. L.; Schwarz, D. E.; Shuh, D. K.; Wagner, G. L.; Wilkerson, M. P.; Wolfsberg, L. E.; Yang, P. Determining Relative f and d Orbital Contributions to M–Cl Covalency in MCl62– (M ¼ Ti, Zr, Hf, U) and UOCl5–Using Cl K-Edge X-Ray Absorption Spectroscopy and Time-Dependent Density Functional Theory. J. Am. Chem. Soc. 2012, 134 (12), 5586–5597. https://doi.org/10.1021/ja2105015. 259. Kozimor, S. A.; Yang, P.; Batista, E. R.; Boland, K. S.; Burns, C. J.; Clark, D. L.; Conradson, S. D.; Martin, R. L.; Wilkerson, M. P.; Wolfsberg, L. E. Trends in Covalency for D- and f-Element Metallocene Dichlorides Identified Using Chlorine K-Edge X-Ray Absorption Spectroscopy and Time-Dependent Density Functional Theory. J. Am. Chem. Soc. 2009, 131 (34), 12125–12136. https://doi.org/10.1021/ja9015759. 260. Bergmann, U.; Glatzel, P. X-Ray Emission Spectroscopy. Photosynth. Res. 2009, 102 (2–3), 255–266. https://doi.org/10.1007/s11120-009-9483-6. 261. Kowalska, J. K.; Lima, F. A.; Pollock, C. J.; Rees, J. A.; DeBeer, S. A Practical Guide to High-Resolution X-Ray Spectroscopic Measurements and Their Applications in Bioinorganic Chemistry. Isr. J. Chem. 2016, 56 (9–10), 803–815. https://doi.org/10.1002/ijch.201600037. 262. Glatzel, P.; Bergmann, U. High Resolution 1s Core Hole X-Ray Spectroscopy in 3d Transition Metal Complexes—Electronic and Structural Information. Coord. Chem. Rev. 2005, 249 (1), 65–95. https://doi.org/10.1016/j.ccr.2004.04.011. 263. Martinie, R. J.; Blaesi, E. J.; Krebs, C.; Bollinger, J. M.; Silakov, A.; Pollock, C. J. Evidence for a Di-m-Oxo Diamond Core in the Mn(IV)/Fe(IV) Activation Intermediate of Ribonucleotide Reductase from Chlamydia Trachomatis. J. Am. Chem. Soc. 2017, 139 (5), 1950–1957. https://doi.org/10.1021/jacs.6b11563. 264. Martinie, R. J.; Blaesi, E. J.; Bollinger, J. M.; Krebs, C.; Finkelstein, K. D.; Pollock, C. J. Two-Color Valence-to-Core X-Ray Emission Spectroscopy Tracks Cofactor Protonation State in a Class I Ribonucleotide Reductase. Angew. Chem. Int. Ed. 2018, 57 (39), 12754–12758. https://doi.org/10.1002/anie.201807366. 265. Martin-Diaconescu, V.; Chacón, K. N.; Delgado-Jaime, M. U.; Sokaras, D.; Weng, T.-C.; DeBeer, S.; Blackburn, N. J. Kb Valence to Core X-Ray Emission Studies of Cu(I) Binding Proteins with Mixed Methionine – Histidine Coordination. Relevance to the Reactivity of the M- and H-Sites of Peptidylglycine Monooxygenase. Inorg. Chem. 2016, 55 (7), 3431–3439. https://doi.org/10.1021/acs.inorgchem.5b02842. 266. Lim, H.; Baker, M. L.; Cowley, R. E.; Kim, S.; Bhadra, M.; Siegler, M. A.; Kroll, T.; Sokaras, D.; Weng, T.-C.; Biswas, D. R.; Dooley, D. M.; Karlin, K. D.; Hedman, B.; Hodgson, K. O.; Solomon, E. I. Kb X-Ray Emission Spectroscopy as a Probe of Cu(I) Sites: Application to the Cu(I) Site in Preprocessed Galactose Oxidase. Inorg. Chem. 2020, 59 (22), 16567–16581. https://doi.org/10.1021/acs.inorgchem.0c02495. 267. Kowalska, J. K.; Hahn, A. W.; Albers, A.; Schiewer, C. E.; Bjornsson, R.; Lima, F. A.; Meyer, F.; DeBeer, S. X-Ray Absorption and Emission Spectroscopic Studies of [L 2 Fe 2 S 2 ] n Model Complexes: Implications for the Experimental Evaluation of Redox States in Iron–Sulfur Clusters. Inorg. Chem. 2016, 55 (9), 4485–4497. https://doi.org/10.1021/acs. inorgchem.6b00295.

Computational Methods in Organometallic Chemistry

209

268. Schwalenstocker, K.; Paudel, J.; Kohn, A. W.; Dong, C.; Van, K. M.; Farquhar, E. R.; Li, F. Cobalt Kb Valence-to-Core X-Ray Emission Spectroscopy: A Study of Low-Spin Octahedral Cobalt(III) Complexes. Dalton Trans. 2016, 45, 14191–14202. 269. Lee, N.; Petrenko, T.; Bergmann, U.; Neese, F.; DeBeer, S. Probing Valence Orbital Composition With Iron Kb X-Ray Emission Spectroscopy. J. Am. Chem. Soc. 2010, 132 (28), 9715–9727. https://doi.org/10.1021/ja101281e. 270. Lassalle-Kaiser, B.; Boron, T. T.; Krewald, V.; Kern, J.; Beckwith, M. A.; Delgado-Jaime, M. U.; Schroeder, H.; Alonso-Mori, R.; Nordlund, D.; Weng, T.-C.; Sokaras, D.; Neese, F.; Bergmann, U.; Yachandra, V. K.; DeBeer, S.; Pecoraro, V. L.; Yano, J. Experimental and Computational X-Ray Emission Spectroscopy as a Direct Probe of Protonation States in Oxo-Bridged Mn IV Dimers Relevant to Redox-Active Metalloproteins. Inorg. Chem. 2013, 52 (22), 12915–12922. https://doi.org/10.1021/ic400821g. 271. Krewald, V.; Lassalle-Kaiser, B.; Boron, T. T.; Pollock, C. J.; Kern, J.; Beckwith, M. A.; Yachandra, V. K.; Pecoraro, V. L.; Yano, J.; Neese, F.; DeBeer, S. The Protonation States of Oxo-Bridged Mn IV Dimers Resolved by Experimental and Computational Mn K Pre-Edge X-Ray Absorption Spectroscopy. Inorg. Chem. 2013, 52 (22), 12904–12914. https:// doi.org/10.1021/ic4008203. 272. Phu, P. N.; Gutierrez, C. E.; Kundu, S.; Sokaras, D.; Kroll, T.; Warren, T. H.; Stieber, S. C. E. Quantification of Ni–N–O Bond Angles and NO Activation by X-Ray Emission Spectroscopy. Inorg. Chem. 2021, 60 (2), 736–744. https://doi.org/10.1021/acs.inorgchem.0c02724. 273. Pollock, C. J.; DeBeer, S. Insights into the Geometric and Electronic Structure of Transition Metal Centers From Valence-to-Core X-Ray Emission Spectroscopy. Acc. Chem. Res. 2015, 48 (11), 2967–2975. https://doi.org/10.1021/acs.accounts.5b00309. 274. Pollock, C. J.; Lancaster, K. M.; Finkelstein, K. D.; DeBeer, S. Study of Iron Dimers Reveals Angular Dependence of Valence-to-Core X-Ray Emission Spectra. Inorg. Chem. 2014, 53 (19), 10378–10385. https://doi.org/10.1021/ic501462y. 275. Pollock, C. J.; Grubel, K.; Holland, P. L.; DeBeer, S. Experimentally Quantifying Small-Molecule Bond Activation Using Valence-to-Core X-Ray Emission Spectroscopy. J. Am. Chem. Soc. 2013, 135 (32), 11803–11808. https://doi.org/10.1021/ja3116247. 276. Cutsail, G. E.; Gagnon, N. L.; Spaeth, A. D.; Tolman, W. B.; DeBeer, S. Valence-to-Core X-Ray Emission Spectroscopy as a Probe of O−O Bond Activation in Cu 2 O 2 Complexes. Angew. Chem. Int. Ed. 2019, 58 (27), 9114–9119. https://doi.org/10.1002/anie.201903749. 277. Lu, T.-T.; Weng, T.-C.; Liaw, W.-F. X-Ray Emission Spectroscopy: A Spectroscopic Measure for the Determination of NO Oxidation States in Fe-NO Complexes. Angew. Chem. Int. Ed. 2014, 53 (43), 11562–11566. https://doi.org/10.1002/anie.201407603. 278. Lancaster, K. M.; Roemelt, M.; Ettenhuber, P.; Hu, Y.; Ribbe, M. W.; Neese, F.; Bergmann, U.; DeBeer, S. X-Ray Emission Spectroscopy Evidences a Central Carbon in the Nitrogenase Iron-Molybdenum Cofactor. Science 2011, 334 (6058), 974–977. https://doi.org/10.1126/science.1206445. 279. Delgado-Jaime, M. U.; Dible, B. R.; Chiang, K. P.; Brennessel, W. W.; Bergmann, U.; Holland, P. L.; DeBeer, S. Identification of a Single Light Atom Within a Multinuclear Metal Cluster Using Valence-to-Core X-Ray Emission Spectroscopy. Inorg. Chem. 2011, 50 (21), 10709–10717. https://doi.org/10.1021/ic201173j. 280. Delgado-Jaime, M. U.; DeBeer, S.; Bauer, M. Valence-to-Core X-Ray Emission Spectroscopy of Iron-Carbonyl Complexes: Implications for the Examination of Catalytic Intermediates. Chem. Eur. J. 2013, 19 (47), 15888–15897. https://doi.org/10.1002/chem.201301913. 281. Eeckhout, S. G.; Safonova, O. V.; Smolentsev, G.; Biasioli, M.; Safonov, V. A.; Vykhodtseva, L. N.; Sikora, M.; Glatzel, P. Cr Local Environment by Valence-to-Core X-Ray Emission Spectroscopy. J. Anal. At. Spectrom. 2009, 24 (2), 215–223. https://doi.org/10.1039/B808345M. 282. Smolentsev, G.; Soldatov, A. V.; Messinger, J.; Merz, K.; Weyhermüller, T.; Bergmann, U.; Pushkar, Y.; Yano, J.; Yachandra, V. K.; Glatzel, P. X-Ray Emission Spectroscopy To Study Ligand Valence Orbitals in Mn Coordination Complexes. J. Am. Chem. Soc. 2009, 131 (36), 13161–13167. https://doi.org/10.1021/ja808526m. 283. MacMillan, S. N.; Walroth, R. C.; Perry, D. M.; Morsing, T. J.; Lancaster, K. M. Ligand-Sensitive but Not -Diagnostic: Evaluating Cr Valence-to-Core X-Ray Emission Spectroscopy as a Probe of Inner-Sphere Coordination. Inorg. Chem. 2015, 54, 205–214. 284. Pollock, C. J.; DeBeer, S. Valence-to-Core X-Ray Emission Spectroscopy: A Sensitive Probe of the Nature of a Bound Ligand. J. Am. Chem. Soc. 2011, 133 (14), 5594–5601. https://doi.org/10.1021/ja200560z. 285. Beckwith, M. A.; Roemelt, M.; Collomb, M.-N.; DuBoc, C.; Weng, T.-C.; Bergmann, U.; Glatzel, P.; Neese, F.; DeBeer, S. Manganese Kb X-Ray Emission Spectroscopy As a Probe of Metal–Ligand Interactions. Inorg. Chem. 2011, 50 (17), 8397–8409. https://doi.org/10.1021/ic200970t. 286. Gennari, M.; Brazzolotto, D.; Pécaut, J.; Cherrier, M. V.; Pollock, C. J.; DeBeer, S.; Retegan, M.; Pantazis, D. A.; Neese, F.; Rouzières, M.; Clérac, R.; Duboc, C. Dioxygen Activation and Catalytic Reduction to Hydrogen Peroxide by a Thiolate-Bridged Dimanganese(II) Complex With a Pendant Thiol. J. Am. Chem. Soc. 2015, 137 (26), 8644–8653. https://doi.org/10.1021/jacs.5b04917. 287. Rees, J. A.; Martin-Diaconescu, V.; Kovacs, J. A.; DeBeer, S. X-Ray Absorption and Emission Study of Dioxygen Activation by a Small-Molecule Manganese Complex. Inorg. Chem. 2015, 54 (13), 6410–6422. https://doi.org/10.1021/acs.inorgchem.5b00699. 288. Schuth, N.; Zaharieva, I.; Chernev, P.; Berggren, G.; Anderlund, M.; Styring, S.; Dau, H.; Haumann, M. Ka X-Ray Emission Spectroscopy on the Photosynthetic Oxygen-Evolving Complex Supports Manganese Oxidation and Water Binding in the S 3 State. Inorg. Chem. 2018, 57 (16), 10424–10430. https://doi.org/10.1021/acs.inorgchem.8b01674. 289. Bergmann, U.; Horne, C. R.; Collins, T. J.; Workman, J. M.; Cramer, S. P. Chemical Dependence of Interatomic X-Ray Transition Energies and Intensities—A Study of Mn Kb00 and Kb2,5 Spectra. Chem. Phys. Lett. 1999, 302 (1), 119–124. https://doi.org/10.1016/S0009-2614(99)00095-0. 290. Maganas, D.; DeBeer, S.; Neese, F. A Restricted Open Configuration Interaction with Singles Method To Calculate Valence-to-Core Resonant X-Ray Emission Spectra: A Case Study. Inorg. Chem. 2017, 56 (19), 11819–11836. https://doi.org/10.1021/acs.inorgchem.7b01810. 291. Tocheva, E. I. Side-On Copper-Nitrosyl Coordination by Nitrite Reductase. Science 2004, 304 (5672), 867–870. https://doi.org/10.1126/science.1095109. 292. Ryu, H.; Park, J.; Kim, H. K.; Park, J. Y.; Kim, S.-T.; Baik, M.-H. Pitfalls in Computational Modeling of Chemical Reactions and How To Avoid Them. Organometallics 2018, 37 (19), 3228–3239. https://doi.org/10.1021/acs.organomet.8b00456. 293. Harvey, J. N.; Himo, F.; Maseras, F.; Perrin, L. Scope and Challenge of Computational Methods for Studying Mechanism and Reactivity in Homogeneous Catalysis. ACS Catal. 2019, 9 (8), 6803–6813. https://doi.org/10.1021/acscatal.9b01537. 294. Tomasi, J.; Mennucci, B.; Cammi, R. Quantum Mechanical Continuum Solvation Models. Chem. Rev. 2005, 105 (8), 2999–3094. https://doi.org/10.1021/cr9904009. 295. Roux, B.; Simonson, T. Implicit Solvent Models. Biophys. Chem. 1999, 78 (1), 1–20. https://doi.org/10.1016/S0301-4622(98)00226-9. 296. Christen, M.; van Gunsteren, W. F. On Searching in, Sampling of, and Dynamically Moving Through Conformational Space of Biomolecular Systems: A Review. J. Comput. Chem. 2008, 29 (2), 157–166. https://doi.org/10.1002/jcc.20725. 297. Vasquez, M.; Nemethy, G.; Scheraga, H. A. Conformational Energy Calculations on Polypeptides and Proteins. Chem. Rev. 1994, 94 (8), 2183–2239. https://doi.org/10.1021/ cr00032a002. 298. Karton, A. A Computational Chemist’s Guide to Accurate Thermochemistry for Organic Molecules. WIREs Comput. Mol. Sci. 2016, 6 (3), 292–310. https://doi.org/10.1002/ wcms.1249. 299. Eisenstein, O.; Ujaque, G.; Lledós, A. What Makes a Good (Computed) Energy Profile?In Topics in Organometallic Chemistry, Springer: Berlin, Heidelberg, 2020;; pp 1–38. https://doi.org/10.1007/3418_2020_57. 300. Bernardi, F.; Bottoni, A.; Olivucci, M.; Robb, M. A.; Schlegel, H. B.; Tonachini, G. Do Supra-Antara Paths Really Exist for 2 + 2 Cycloaddition Reactions? Analytical Computation of the MC-SCF Hessians for Transition States of Ethylene with Ethylene, Singlet Oxygen, and Ketene. J. Am. Chem. Soc. 1988, 110 (18), 5993–5995. https://doi.org/10.1021/ ja00226a011. 301. Krylov, A. I.; Sherrill, C. D.; Byrd, E. F. C.; Head-Gordon, M. Size-Consistent Wave Functions for Nondynamical Correlation Energy: The Valence Active Space Optimized Orbital Coupled-Cluster Doubles Model. J. Chem. Phys. 1998, 109 (24), 10669–10678. https://doi.org/10.1063/1.477764. 302. Szalay, P. G.; Müller, T.; Gidofalvi, G.; Lischka, H.; Shepard, R. Multiconfiguration Self-Consistent Field and Multireference Configuration Interaction Methods and Applications. Chem. Rev. 2012, 112 (1), 108–181. https://doi.org/10.1021/cr200137a. 303. Mato, J.; Gordon, M. S. Analytic Gradients for the Spin-Flip ORMAS-CI Method: Optimizing Minima, Saddle Points, and Conical Intersections. J. Phys. Chem. A 2019, 123 (6), 1260–1272. https://doi.org/10.1021/acs.jpca.8b11569.

210

Computational Methods in Organometallic Chemistry

304. Schmidt, M. W.; Gordon, M. S. The Construction and Interpretation of Mcscf Wavefunctions. Annu. Rev. Phys. Chem. 1998, 49 (1), 233–266. https://doi.org/10.1146/annurev. physchem.49.1.233. 305. Odoh, S. O.; Manni, G. L.; Carlson, R. K.; Truhlar, D. G.; Gagliardi, L. Separated-Pair Approximation and Separated-Pair Pair-Density Functional Theory. Chem. Sci. 2016, 7 (3), 2399–2413. https://doi.org/10.1039/C5SC03321G. 306. Ghosh, S.; Cramer, C. J.; Truhlar, D. G.; Gagliardi, L. Generalized-Active-Space Pair-Density Functional Theory: An Efficient Method to Study Large, Strongly Correlated, Conjugated Systems. Chem. Sci. 2017, 8 (4), 2741–2750. https://doi.org/10.1039/C6SC05036K. 307. Andersson, K.; Malmqvist, P. A.; Roos, B. O.; Sadlej, A. J.; Wolinski, K. Second-Order Perturbation Theory With a CASSCF Reference Function. J. Phys. Chem. 1990, 94 (14), 5483–5488. https://doi.org/10.1021/j100377a012. 308. Hirao, K. Multireference Møller—Plesset Perturbation Theory for High-Spin Open-Shell Systems. Chem. Phys. Lett. 1992, 196 (5), 397–403. https://doi.org/10.1016/00092614(92)85710-R. 309. Hirao, K. State-Specific Multireference Møller—Plesset Perturbation Treatment for Singlet and Triplet Excited States, Ionized States and Electron Attached States of H2O. Chem. Phys. Lett. 1993, 201 (1), 59–66. https://doi.org/10.1016/0009-2614(93)85034-L. 310. Shepard, R. Geometrical Energy Derivative Evaluation With MRCI Wave Functions. Int. J. Quantum Chem. 1987, 31 (1), 33–44. https://doi.org/10.1002/qua.560310105. 311. Sand, A. M.; Kidder, K. M.; Truhlar, D. G.; Gagliardi, L. Calculation of Chemical Reaction Barrier Heights by Multiconfiguration Pair-Density Functional Theory With Correlated Participating Orbitals. J. Phys. Chem. A 2019, 123 (45), 9809–9817. https://doi.org/10.1021/acs.jpca.9b08134. 312. Kenion, R. L.; Ananth, N. Direct Simulation of Electron Transfer in the Cobalt Hexammine(II/III) Self-Exchange Reaction. Phys. Chem. Chem. Phys. 2016, 18 (37), 26117–26124. https://doi.org/10.1039/C6CP04882J. 313. García-López, D.; Pavlovic, L.; Hopmann, K. H. To Bind or Not to Bind: Mechanistic Insights into C–CO2 Bond Formation With Late Transition Metals. Organometallics 2020, 39 (8), 1339–1347. https://doi.org/10.1021/acs.organomet.0c00090. 314. Ditchfield, R.; Hehre, W. J.; Pople, J. A. Self-Consistent Molecular-Orbital Methods. IX. An Extended Gaussian-Type Basis for Molecular-Orbital Studies of Organic Molecules. J. Chem. Phys. 1971, 54 (2), 724–728. https://doi.org/10.1063/1.1674902. 315. Hehre, W. J.; Ditchfield, R.; Pople, J. A. Self—Consistent Molecular Orbital Methods. XII. Further Extensions of Gaussian—Type Basis Sets for Use in Molecular Orbital Studies of Organic Molecules. J. Chem. Phys. 1972, 56 (5), 2257–2261. https://doi.org/10.1063/1.1677527. 316. Dill, J. D.; Pople, J. A. Self-Consistent Molecular Orbital Methods. XV. Extended Gaussian-Type Basis Sets for Lithium, Beryllium, and Boron. J. Chem. Phys. 1975, 62 (7), 2921–2923. https://doi.org/10.1063/1.430801. 317. Francl, M. M.; Pietro, W. J.; Hehre, W. J.; Binkley, J. S.; Gordon, M. S.; DeFrees, D. J.; Pople, J. A. Self-Consistent Molecular Orbital Methods. XXIII. A Polarization-type Basis Set for Second-Row Elements. J. Chem. Phys. 1982, 77 (7), 3654–3665. https://doi.org/10.1063/1.444267. 318. Grimme, S.; Antony, J.; Ehrlich, S.; Krieg, H. A Consistent and Accurate Ab Initio Parametrization of Density Functional Dispersion Correction (DFT-D) for the 94 Elements H-Pu. J. Chem. Phys. 2010, 132 (15), 154104. https://doi.org/10.1063/1.3382344. 319. Grimme, S.; Ehrlich, S.; Goerigk, L. Effect of the Damping Function in Dispersion Corrected Density Functional Theory. J. Comput. Chem. 2011, 32 (7), 1456–1465. https://doi. org/10.1002/jcc.21759. 320. Scalmani, G.; Frisch, M. J. Continuous Surface Charge Polarizable Continuum Models of Solvation. I. General Formalism. J. Chem. Phys. 2010, 132 (11), 114110. https://doi. org/10.1063/1.3359469. 321. Jindal, G.; Kisan, H. K.; Sunoj, R. B. Mechanistic Insights on Cooperative Catalysis through Computational Quantum Chemical Methods. ACS Catal. 2015, 5 (2), 480–503. https://doi.org/10.1021/cs501688y. 322. Peltzer, R. M.; Gauss, J.; Eisenstein, O.; Cascella, M. The Grignard Reaction—Unraveling a Chemical Puzzle. J. Am. Chem. Soc. 2020, 142 (6), 2984–2994. https://doi.org/ 10.1021/jacs.9b11829. 323. Ahn, S.; Hong, M.; Sundararajan, M.; Ess, D. H.; Baik, M.-H. Design and Optimization of Catalysts Based on Mechanistic Insights Derived from Quantum Chemical Reaction Modeling. Chem. Rev. 2019, 52, 6509–6560. 324. Grignard, V. C. Sur Quelques Nouvelles Combinaisons Organométalliques Du Magnésium et Leur Application à Des Synthèses Daalcools et d’hydrocarbures. C. R. Hebd. Seances Acad. Sci. 1900, 130, 1322–1324. 325. Seyferth, D. The Grignard Reagents. Organometallics 2009, 28 (6), 1598–1605. https://doi.org/10.1021/om900088z. 326. Peltzer, R. M.; Eisenstein, O.; Nova, A.; Cascella, M. How Solvent Dynamics Controls the Schlenk Equilibrium of Grignard Reagents: A Computational Study of CH3MgCl in Tetrahydrofuran. J. Phys. Chem. B 2017, 121 (16), 4226–4237. https://doi.org/10.1021/acs.jpcb.7b02716. 327. Walker, F. W.; Ashby, E. C. Composition of Grignard Compounds. VI. Nature of Association in Tetrahydrofuran and Diethyl Ether Solutions. J. Am. Chem. Soc. 1969, 91 (14), 3845–3850. https://doi.org/10.1021/ja01042a027.

1.08

f-Element Organometallic Single-Molecule Magnets

Richard A Layfield, Christopher GT Price, and Siobhan R Temple, Department of Chemistry, School of Life Sciences, University of Sussex, Brighton, United Kingdom © 2022 Elsevier Ltd. All rights reserved.

1.08.1 Introduction 1.08.1.1 Single-molecule magnetism 1.08.2 Single-molecule magnetism in lanthanide organometallics 1.08.2.1 Lanthanide metallocene single-molecule magnets 1.08.2.1.1 SMMs based on [Cp2Ln(m-X)]n metallocene units 1.08.2.1.2 Lanthanide half-sandwich complexes as SMMs 1.08.2.1.3 Lanthanide metallocene SMMs with radical bridging ligands 1.08.2.1.4 Cationic dysprosium metallocene SMMs [(CpR)2Dy]+ 1.08.2.2 Lanthanide single-molecule magnets based on cyclooctatetraene ligands 1.08.2.3 Organometallic lanthanide SMMs containing 4-, 6-, or 7-membered rings 1.08.2.3.1 Lanthanide SMMs with Z4-cyclobutadienyl ligands 1.08.2.3.2 Lanthanide SMMs with Z6-arene or Z7-cycloheptatrienyl ligands 1.08.2.4 Lanthanide organometallic SMMs based on s-bonded ligands 1.08.2.5 Single-molecule magnetism in actinide organometallics 1.08.3 Conclusions and outlook Acknowledgment References

1.08.1

211 212 215 215 215 221 223 226 230 236 236 238 240 243 245 245 246

Introduction

Organometallic chemistry has a long and distinguished history dating back over two centuries.1 As earlier incarnations of Comprehensive Organometallic Chemistry and even a cursory reading of the organometallic literature reveal, applications of compounds containing metal-carbon bonds are most readily associated with a bewildering variety of catalytic and stoichiometric reactions. Areas previously regarded as emerging, such as medicinal organometallic chemistry2 and organometallic polymers,3 have become firmly established research themes in their own right. The predominance of 18- and 16-electron configurations has meant that the closed-shell nature of most organotransition metal compounds is taken as ‘normal’, whereas open-shell or paramagnetic organometallics tend to be regarded as the exceptions. On the other hand, the established discipline of molecular magnetism (or magnetochemistry) typically conjures images of unpaired electrons and spin Hamiltonians associated with classical, Werner-type coordination compounds4 or main group radicals.5 While magnetism in organometallic compounds is not a revolutionary concept, as a field of study it has evolved along separate lines to conventional molecular magnetism. Organometallic magnets have very rarely been subjected to the arsenal of analytical, spectroscopic and theoretical techniques used widely in molecular magnetism to establish quantitative explanations of the complex phenomena arising from unpaired electron spin. Despite this, unusual magnetic phenomena were discovered in the early days of post-ferrocene organometallic chemistry, including the peculiar family of divalent manganocenes,6 the spin configurations of which are remarkably sensitive to the ligand substituents, leading to spin-crossover phenomena more typical of classical coordination compounds. The famous charge-transfer salts based on ferrocene electron-donors in combination with stable radical anions derived from, e.g., tetracyanoethylene or tetrathiafulvalene, are among the few organometallic compounds to show genuine ferromagnetic properties.7 Despite the early pioneering work in organometallic magnetism, the development of organometallic compounds as a mainstream branch of molecular magnetism is much more recent. Over the last decade in particular, the metal-carbon bond as a design tool in molecular magnetism has grown in importance owing to the discovery of the first organometallic single-molecule magnets (SMMs).8 A single, textbook definition of an SMM has not yet been agreed, although these materials have been variously classified according to their ability to retain magnetization in the absence of an applied magnetic field below a certain temperature characteristic of the molecule, or in terms of the relatively slow rate at which the magnetization relaxes, or in terms of an effective energy barrier required to flip the molecular magnetic dipole. This chapter is dedicated to the most important findings in this vibrant field, with a strong emphasis on lanthanide organometallic compounds and with notable contributions from actinide organometallic chemistry. Another important set of compounds with fascinating magnetic and spectroscopic properties are multiconfigurational sandwich complexes of cerium and ytterbium. Oxidation state ambiguity has become the hallmark of formally tetravalent cerocene [Ce(Z8-COT)2] (COT ¼ cyclooctatetraene)9–11 and formally divalent decamethylytterbocene [Yb(Z5-Cp )2] (Cp ¼ C5Me5),12,13 a phenomenon that has provided a rich source of intrigue surrounding the admixture of the trivalent metal oxidation state and a ligand radical into the electronic and magnetic ground states.14

Comprehensive Organometallic Chemistry IV

https://doi.org/10.1016/B978-0-12-820206-7.00070-6

211

212

1.08.1.1

f-Element Organometallic Single-Molecule Magnets

Single-molecule magnetism

The phenomenon of single-molecule magnetism has been observed in hundreds of coordination and organometallic compounds containing metal ions with certain key physical properties. The properties and conditions needed to observe SMM behavior include magnetic anisotropy and unpaired spin that can be magnetized (or oriented) in an applied magnetic field (H in Oersted units). When the field is switched off (H ¼ 0), the SMM molecules retain at least some magnetization (M) for a period of time before this quantity decays to zero through one or more spin-lattice relaxation processes.15 Magnetization in SMMs is retained only below a certain temperature characteristic of the molecule, and the relaxation time also typically shortens with increasing temperature. SMMs are therefore distinct from simple Curie paramagnets, where the magnetization is entirely dependent on the application of an external magnetic field, regardless of the temperature. At the time of writing, all SMMs worthy of the name display magnetic relaxation well within the cryogenic temperature regime (i.e. below 123 K), and in most cases not above liquid-helium temperatures. One of the main challenges confronting researchers in the field is, therefore, to develop SMMs that function at higher temperatures. Since the discovery of the first SMMs in the early 1990s,16,17 applications have been proposed although not yet realized. The main advances have been fundamental in nature. However, the rate of progress is accelerating in a manner that gives some hope that applications of SMMs in magnetic information storage or as the basis of molecular spintronic devices may not be such an unrealistic aspiration. Review articles focusing on the prospects for functional molecular magnetic materials are available.18,19 This chapter will primarily explore the underlying concepts behind f-element SMMs and how organometallic chemistry has been able to advance the field in a manner that complements the important progress made with Werner-type coordination compounds. SMMs based on 3d organometallic compounds are not covered in this chapter, however the reader can be directed to the key articles describing the magnetic properties of low-coordinate iron,20–24 cobalt,25–27 and nickel.28 The key ingredients required to observe SMM behavior are spin-orbit coupling (SOC) and magnetic anisotropy. The resulting aspherical electron density of a particular metal ion is what leads to a preferred orientation of the molecular magnetic moment, along which it is ‘easy’ to magnetize the molecule: reassuringly, this is referred to as the easy axis of magnetization. The number of unpaired electrons in the molecule (leading to the total spin, Stot) is not the primary consideration in determining SMM behavior. Since SOC is the important property, by far the most popular metal ions found in contemporary SMMs are trivalent lanthanide ions (Ln3+).29–35 The 4fn configurations of trivalent lanthanides experience weak crystal field effects, meaning that the orbital contribution to the magnetism is essentially unquenched, in contrast to transition metal and actinide ions. Indeed, effective magnetic moments (meff) for Ln3+ ions containing unpaired spin often deviate substantially from the spin-only value except in the case of isotropic Gd3+ (4f7).36 It is important to emphasize that not every anisotropic lanthanide is suited to SMM behavior. A select group from terbium to erbium accounts for almost all lanthanide SMMs, with dysprosium dominating the landscape. The reasons for the prevalence of these lanthanides in SMMs are simultaneously simple and complicated. Space constraints preclude detailed discussion of the underlying physics, which would stray too far from the remit of Comprehensive Organometallic Chemistry, and this topic has been detailed in reviews and monographs.4,36,37 In short, these heavier lanthanide ions have highly anisotropic 4fn electron density and the SOC is very strong. The magnitude of the SOC is reflected in the values of the total SOC quantum number J in the 2S+1LJ ground multiplets of Tb3+, Dy3+, Ho3+ and Er3+ (Table 1). The ground multiplets of the Ln3+ ions interact with the crystal field generated by the ligands. Alluded to above, the crystal field splitting is weak but non-negligible. The nature of the interaction has been debated in the context of SMMs and, while the overall metal-ligand bonding is predominantly electrostatic in nature, an appreciable covalent contribution is responsible for the splitting of the ground multiplet into a series of crystal field states,38,39 of which there are 2J + 1 for each Ln3+ ion. The covalency is thought to involve donation of electron density from the ligands to the lanthanide 5d orbitals. Softer, more polarizable ligands such as those commonly employed in organometallic chemistry can, in principle, provide greater covalent contributions than ligands containing the traditional hard donor atoms. Both Dy3+ (4f9) and Er3+ (4f11) are Kramers ions (non-integer spin), one consequence of which, according to Kramers’ theorem, is that the crystal field states always occur as a series of doublets. The two components of each Kramers doublet (KD) are then connected, mathematically, via time-reversal symmetry.40 In contrast, Tb3+ (4f8) and Ho3+ (4f10) are non-Kramers ion (integer spin), and the resulting doublets only occur if the ion occupies an environment with strict axial symmetry. In an ideal coordination geometry/symmetry, each crystal field state is defined by a definite value of the angular momentum projection  MJ. The nature of the ideal geometry depends on the shape of the anisotropic 4f electron density associated with a Table 1

Selected properties of heavier lanthanide trivalent cations.

Ln3+

4fn

Stot

2S+1

gJ

meff/mB

Gd3+ Tb3+ Dy3+ Ho3+ Er3+ Tm3+ Yb3+

7 8 9 10 11 12 13

7/2 3 5/2 2 3/2 1 1/2

8

2 3/2 4/3 5/4 6/5 7/6 8/7

7.94 9.72 10.65 10.61 9.58 7.56 4.54

LJ

S7/2 F6 6 H15/2 5 I8 4 I5/2 3 H6 2 F7/2 7

f-Element Organometallic Single-Molecule Magnets

213

particular ion. In the case of Tb3+, Dy3+ and Ho3+ the electron density is an oblate spheroidal shape (compressed along the z-direction), whereas in Er3+ it is a prolate spheroidal shape (stretched along the z-direction) (Fig. 1).41,42 For oblate ions, the ideal geometry for SMM behavior is linear and two-coordinate since the crystal field arising from such an arrangement is purely axial, which enhances the single-ion anisotropy.43,44 Conversely, for prolate ions the ideal geometry is planar with no axial contribution, such as trigonal planar. The oblate model is exemplified with Dy3+ in Fig. 2 for a hypothetical complex cation [DyZ2]+, where Z is a monoanionic ligand and the Z–Dy–Z angle is 180 . In this ideal picture, the first key SMM metric is introduced. A flip of the magnetic dipole from the MJ ¼ + 15/2 state to the MJ ¼ −15/2 state, in quantized steps via the highest-lying KD, defines the theoretical maximum energy barrier to relaxation of the magnetization, or Utheor. Part of the reason why dysprosium is so prevalent in SMMs is that some sort of doublet structure, derived from the ideal crystal field splitting in Fig. 2, persists even when the geometry deviates from perfectly axial. Indeed, perfectly axial two-coordinate Dy3+ SMMs are, at the time of writing, unknown. A consequence of a bent geometry and/or an equatorial component of the crystal field is mixing of the MJ states, leading to thermal relaxation via an under-barrier shortcut, such that the measured effective energy barrier to reversal of the magnetization, Ueff, is always less than Utheor for SMMs containing lanthanides with multilevel crystal field splitting. An ab initio theoretical study of the unknown but chemically not entirely unrealistic complex [(Dy(CAAC)2]3+, where CAAC denotes a cyclic alkyl amino carbene with a 2,6-diisopropylphenyl substituent on the amino nitrogen atom, highlighted many of the main points associated with the electronic structure of Dy3+ ions in a linear, two-coordinate environment.43 In the crystal field splitting diagram for this carbene complex, all eight KDs are shown in Fig. 3. The most probable relaxation pathway consists of a series of spin-phonon transitions represented by the red arrows, meaning that relaxation in [(Dy(CAAC)2]3+ begins in the ground KD with MJ ¼ +15/2 and proceeds in a stepwise manner up to the fifth KD, where a barrier-crossing transition occurs. Subsequently, the system relaxes down the other side of the barrier, eventually reaching the time-reversed component of the ground KD with MJ ¼ − 15/2. Overall, in this model system, the effective energy barrier is Ueff ¼ 1750 cm−1. Furthermore, since the quantum tunneling of the magnetization (QTM) processes (green arrows in Fig. 3) are improbable in the lowest-lying KDs, this linear two-coordinate SMM would also be expected to show magnetic hysteresis loops with appreciable coercive fields at low temperatures. It is noteworthy that, even in this model system, activated relaxation still does not occur via the eighth and highest-lying KD. Therefore, in real systems, where the geometry deviates from perfectly axial, the measured energy barrier Ueff invariably reflects a relaxation mechanism in which a shortcut under the maximum possible barrier occurs. Despite this, a limited number of theoretical

Fig. 1 The 4fn electron density in Ln3+ ions. Adapted with permission from reference Jiang, S.-D.; Qin S.-X. Inorg. Chem. Front. 2015, 2, 613–619.

Fig. 2 Schematic representation of possible magnetic relaxation pathways in the four (of eight) lowest-lying Kramers doublets of Dy3+ in a strictly linear, two-coordinate complex [DyZ2]+.

214

f-Element Organometallic Single-Molecule Magnets

2000

E (cm–1)

1600

3.3

1200 –11/2

DippN

800 400

–13/2

–15/2 –10

+9/2

–3

10

0.23 x 10–3

+11/2

–4

0.39

2.1 0

0.24 x 10

0.80 x

2.8

+7/2 –1

–9/2

NDipp

Dy

0.83

–7/2 3.0

x 10

0.26 x 10–4

+13/2

–6

0.14

–8

x 10

–6

0.53 x 10–6 –4

–2

0

2

+15/2 4

6

8

10

M (mB) Fig. 3 Molecular structure of hypothetical, linear [Dy(CAAC)2]3+ and splitting of the crystal field levels into Kramers doublets. The dashed red line on the molecule indicates the easy axis of magnetization. The solid red arrows on the energy spectrum denote the most probable relaxation pathway, the green arrows indicate quantum tunneling of the magnetization, and blue arrows denote other possible Orbach relaxation processes. The numerical values above the arrows are the average transition dipole moments. Adapted with permission from reference Ungur, L.; Chibotaru, L. F. Inorg. Chem. 2016, 55, 10043–10056.

studies have indicated that a reasonable degree of bending in two-coordinate dysprosium SMMs can be tolerated without impacting on the magnitude of Ueff,43,44 a useful result indicating that perfect axiality is not essential for a large effective energy barrier. Discussed in detail in Section 1.08.2.1.4, bent pseudo-two coordinate dysprosium metallocenes represent the best-performing SMMs to date, with the current ‘record’ Ueff being 1541 cm−1 for [(Ci5Pr5)Dy(Cp )][B(C6F5)4].45 Other SMMs are known in which the crystal field has dominant axial character but with a non-negligible equatorial contribution, such as that found in pentagonal bipyramidal [Dy(OtBu)2(py)5][BPh4] (py ¼ pyridine), which has Ueff ¼ 1261(1) cm−1.46 It has also been shown that the absence of an equatorial crystal field is vital for observing magnetic hysteresis with remnance and coercivity. Indeed, while large Ueff values can be observed for SMMs with non-negligible equatorial crystal fields, these systems tend to have relatively poor hysteresis properties. The result is a low magnetic blocking temperature (TB), usually below 10 K. Leading SMMs without appreciable equatorial crystal fields, which just happen to be dysprosium metallocenes, can have blocking temperatures in the region 60–80 K.45,47–49 As depicted in Fig. 3, the red-arrows transitions denote Orbach processes. In any given SMM, multiple Orbach processes could occur between different states, each with a finite probability, with the most probable process typically corresponding to Ueff. The process is initiated when a phonon of an appropriate energy is absorbed from the lattice, promoting the system to a higher-lying KD, which then relaxes to the other component of the ground KD and emits a second phonon back into the lattice. Several other relaxation mechanisms of the spin-lattice type are also available to SMMs. In the so-called direct process, a transition from one spin orientation to another occurs when the energy difference between those two states corresponds to the energy of a lattice vibration. The direct transition does not proceed via a higher-lying state and emits a phonon into the lattice. In the two-phonon Raman process, one phonon is absorbed at the same time as another is emitted, with the relaxation proceeding via a so-called virtual excited state, which is effectively a superposition of different states arising from multiple lattice vibrations. Another important SMM relaxation process involves through-barrier transitions via resonant QTM. The QTM process can either involve an excited KD, resulting in thermally assisted QTM, or it can occur in the ground KD. Overall, the dependence of the magnetic relaxation time (t) can be expressed as a sum of the individual processes, according to Eq. (1): −1 t −1 ¼ t0−1 e −Ueff =kB T + CT n + tQTM

(1)

In Eq. (1), the first term is the Orbach term, which accounts for an Arrhenius-type dependence of t on temperature and allows reversal of the magnetization to be likened to an activation energy.15 The pre-exponential factor t0 is the attempt frequency and can be taken as a measure of how likely this process is to contribute to the relaxation, with larger values implying higher probability. In the Raman term CTn, C is the Raman coefficient and n is the Raman exponent, taking values in the range n ¼ 2–9 depending on the molecule. The rate of the QTM process is t−1 QTM. An additional term not shown in Eq. (1), accounting for direct relaxation, is AB4T, where A is a measure of how many phonons are available and B is the magnetic field (hence direct relaxation is not important in zero applied magnetic field). Values of t are normally determined from alternating current (AC) magnetic susceptibility measurements.50 The temperature dependence of t, usually depicted as ln t versus T−1, produces a graph that can be fitted using some or all of the terms in equation 1. A detailed description of the AC magnetic susceptibility measurements commonly used to characterize SMMs is beyond the scope of this chapter, suffice to state that the most widely used approach determines the frequency (n) dependence of the imaginary component of the susceptibility (w”) at various temperatures, or the temperature-dependence of w” at various frequencies in the range 0.1 Hz up to 1500 or 10,000 Hz. Generally, the w”(n) data at a given temperature rises to a maximum value and then decreases towards zero, with the maximum at each temperature being used to extract t using the relationship t ¼ 1/(2pn).

f-Element Organometallic Single-Molecule Magnets

215

The complexity of magnetic relaxation in SMMs is underscored by the frequent observation of multiple processes in the same system, with the dominance of one process over the others depending on temperature. Orbach processes tend to occur at higher temperatures, QTM is typically a low-temperature process, and the Raman process can either dominate the relaxation at higher temperatures or occur in an intermediate temperature regime between Orbach and QTM processes. Raman-only relaxation is also known. The goal of the synthetic (organometallic) chemist is to exert some sort of control over the electronic structure such that the Orbach process is dominant with the largest possible Ueff energy barrier. Organometallic chemistry, especially that of the lanthanides, has contributed significantly towards achieving this goal.

1.08.2

Single-molecule magnetism in lanthanide organometallics

In keeping with the general trend in studies of SMMs, those based on lanthanides form by far the largest subset of organometallic derivatives.51–54 For organisational purposes, organolanthanide SMMs are further divided in this chapter by ligand type, although many heteroleptic SMMs containing different types of organometallic ligand have been reported. Sandwich complexes are prevalent in the field, hence the coverage spans ring sizes ranging from Z4-cyclobutadienyl to Z9-cyclononatetraenyl, both of which are relatively rare. The most common organometallic ligands in lanthanide SMMs are Z5-cyclopentadienyl and then Z8-cyclooctatetraenyl. Unless otherwise stated, all Zn-bonded ligands described in this chapter are assumed to adopt their maximum hapticity owing to the relatively large radii of Ln3+ ions. Alkyl and aryl ligands in lanthanide SMMs are also uncommon in the SMM arena, as are ‘pseudo-organometallic’ ligands such as amido or hydrido, and are considered at various point throughout the chapter. For ease of comparison, all Ueff values are stated in units of wavenumbers even if originally reported in kelvin.

1.08.2.1 1.08.2.1.1

Lanthanide metallocene single-molecule magnets SMMs based on [Cp2Ln(m-X)]n metallocene units

The first organometallic SMM of any kind was reported in 2010. The deprotonation of benzotriazole (btaH) by Cp3Dy produced the dimer [Cp2Dy(m-bta)]2, in which both [bta]− ligands adopt an unsymmetrical k2:k1 bridging mode between the Dy3+ centers (Fig. 4).55 Below 12 K in zero applied DC field, the imaginary component of the AC susceptibility is temperature dependent, and maxima in w” were found below 8 K. The temperature dependence of t for this SMM is approximately linear in the region 6.3–10 K

Fig. 4 Structures of dysprosium metallocene SMMs.

216

f-Element Organometallic Single-Molecule Magnets

before switching to a temperature-independent regime at lower temperatures, indicative of QTM. The energy barrier for the thermally activated Orbach process was determined to be 22  1 cm−1 with a pre-exponential factor of t0 ¼ 4.5  10−7 s. The barrier increases to 27  0.4 cm−1 in an applied DC magnetic field of 1000 Oe. A study of [Cp2Dy(m-bta)]2 using density functional theory (DFT) yielded Mulliken population parameters and Mayer bond orders that suggested very weak interactions between the Dy3+ ions, implying that the SMM properties are single-ion in nature despite the dimetallic nature of the molecule. Although perhaps not fully appreciated at the time of the original report, the weak interactions between dysprosium and the equatorial [bta]− ligands, complemented by relatively strong interactions with the axial [Cp]− ligands, are the important criteria for observing SMM behavior in dysprosium metallocenes. SMM behavior was observed in two, co-crystallized polymorphs of [Cp2Dy(m-Cl)]n, where n ¼ 2 or 1, and the solvated complex [Cp2Dy(m-Cl)(THF)]2 (Fig. 4).56 The w”(T ) data for [Cp2Dy(m-Cl)]n, which forms as a 3:1 mixture of dimer and polymer, consists of two sets of peaks below 40 K. Separate analysis of the two sets of peaks produced energy barriers of Ueff ¼ 26.3  0.7 cm−1 and 68  1.2 cm−1 for the dimer and polymer, respectively, with t0 ¼ 1.4  10−6 s and t0 ¼ 2.8  10−6 s. A barrier of Ueff ¼ 34  0.5 cm−1 with t0 ¼ 4.0  10−7 s was determined for [Cp2Dy(m-Cl)(THF)]2. At lower temperatures, magnetic relaxation in the chloride-bridged metallocene dimers is dominated by QTM processes, but in the chloride bridged polymer the relaxation time remains strongly temperature-dependent, even in applied fields up to 5000 Oe. Indeed, the relaxation time at 8 K is 0.15 ms in [Cp2Dy(m-Cl)]2 but 78.6 ms in [Cp2Dy(m-Cl)]1, i.e. 500 times longer as a consequence of the chain-like nature of the polymer. When replacing hard N- or Cl-donor ligands with the softer thiolate donor in the dimer [(CpMe)2Dy(m-SSiPh3)]2 (Figs. 4 and 5, CpMe ¼ methylcyclopentadienyl), a substantial increase in the energy barrier to 133  3.5 cm−1 was observed.57 This barrier was extracted from maxima in the w”(T ) in the region 20–32 K (Fig. 5), where the dependence of ln t on T−1 is linear. At lower temperatures, the relaxation time does not become temperature independent, implying that QTM is inefficient in this SMM and that Raman processes play a role in the relaxation. Quantitative insight into the electronic structure of [Cp2Dy(m-bta)]2, [Cp2Dy(m-Cl)]n, [Cp2Dy(m-Cl)(THF)]2 and [(CpMe)2Dy (m-SSiPh3)]2 was obtained using multireference ab initio calculations in a study that provided the first explanation of the origins of the SMM properties of dysprosium metallocene SMMs.57 The calculations revealed that the crystal field interaction is dominated by the cyclopentadienyl ligands, which adopt axial positions with respect to Dy3+, and the heteroatom donor ligands that bridge between the metals provide a competing equatorial crystal field. This situation is illustrated by considering the easy axis of magnetization in the ground KD for [(CpMe)2Dy(m-SSiPh3)]2, which is clearly oriented towards the [CpMe]− ligands (Fig. 5). A similar picture emerges for the other metallocene SMMs considered in this study. The ab initio calculations also provided a wealth of additional information regarding mechanistic aspects of the magnetic relaxation. For example, in [(CpMe)2Dy(m-SSiPh3)]2 the energy separation between the ground and the first-excited KD was calculated to be 113 cm−1, whereas the gap to the second-excited KD is 283 cm−1. Consequently, the Orbach relaxation in the thiolate-bridged metallocene SMM is likely to proceed via the first-excited KD and not via higher-lying KDs. Note that the slight discrepancy between the calculated and measured barriers is not unusual for an SMM. This may be a consequence of, for example, experimental error in the susceptibility measurement or sample mass, too few or too many data points used in the fitting process, or limitations with the calculations, including the omission of electron correlation effects outside of the 4f orbital manifold. The calculations also unravel the nature of the magnetic interactions between the Dy3+ ions, which can be difficult to model experimentally in a reliable manner owing to the number of parameters required with anisotropic exchange. The consistent picture is that the total exchange, which is a sum of the superexchange and dipolar contributions, i.e. Jtot ¼ Jexch + Jdip, is typically dominated by the dipolar part in SMMs of this type. In general,

Fig. 5 Left: Molecular structure of [(CpMe)2Dy(m-SSiPh3)]2 with the easy axis of magnetization in the ground KDs indicated by dashed red lines. Right: temperature dependence of the imaginary component of the AC magnetic susceptibility (w”) at various frequencies (n).

f-Element Organometallic Single-Molecule Magnets

217

the overall interaction in dimetallic compounds with the general formula [Cp2Dy(m-X)]2 is antiferromagnetic; however, it is also usually weak, such that the SMM properties tend to be dominated by single-ion effects. The situation changes dramatically when the bridging ligands are radicals (see Section 1.08.2.1.3). With this emerging picture of SMM behavior in dysprosium metallocenes, it is possible to account for their magnetic hysteresis properties, which are generally poor when compared to traditional ferromagnetic materials. In a ferromagnet, saturation magnetization (Msat) is eventually reached by applying a magnetic field (H/Oe) of increasing strength. When the field is returned to zero, remanent magnetization (Mr) is observed. Driving the field in the opposite direction is required to return the magnetization to zero, which occurs at the coercive field (Hc). Continuing to sweep the field in a series of forward and reverse directions produces hysteresis loops, leading to magnetic bistability. The appearance or width of the loops depends strongly on temperature and the field sweep rate (in Oe s−1), with loops typically being widest at low temperatures and with fast sweeping. The hysteresis properties of, e.g. [(CpMe)2Dy(m-SSiPh3)]2, such as they are, consist of very narrow loops even at 1.8 K, with negligible remanence and coercivity. These so-called waist-restricted hysteresis loops imply fast magnetic relaxation, which is a consequence of the equatorial crystal field provided by the heteroatom donor ligands, with the exchange coupling also thought to play a role by facilitating QTM processes. The easy axis of magnetization was determined both experimentally and theoretically for the series of centrosymmetric metallocene dimers [Cp’2Dy(m-X)]2, in which Cp’ is trimethylsilylcyclopentadienyl and X ¼ CH3, Cl, Br and I (Fig. 4). SMM properties were observed for the halide-bridged dimers in zero DC field, with energy barriers of Ueff ¼ 106, 139 and 341 cm−1 for the chloride-, bromide- and iodide-ligated versions, respectively.58 Relaxation via QTM was not observed below 10 K. The crystal symmetry of these compounds allowed the easy axes of magnetization to be determined using angular-resolved magnetometry measurements on oriented single crystals. This study confirmed that the easy axes are oriented almost perpendicular (84.4–86.0 ) relative to the {Dy2X2} rings and towards the Cp’ ligands. The experimental results were fully consistent with an ab initio theoretical study, which revealed that the dominant thermal relaxation process occurs via the first-excited KDs in these compounds. The methyl-bridged complex [Cp’2Dy(m-CH3)]2 showed no slow relaxation properties in the absence of an applied field, reflecting the relatively strong donor properties of this ligand in an equatorial position. Bulk AC susceptibility measurements on the monometallic compounds [(Cp )2Dy(X)(THF)] (X ¼ Cl, Br, I) (Fig. 4) also highlight the correlation between SMM behavior and the interaction with the equatorial ligands, with Ueff increasing in the order 112, 163 and 419 cm−1 when varying the halide from chloride to bromide to iodide.59 In the case of [(Cp )2Dy(I)(THF)], the Orbach relaxation is thought to proceed via the excited states possibly up to the third-excited KD, whereas [(Cp )2Dy(X)(THF)] with X ¼ Cl or Br should relax via the first-excited KD. In a complementary manner, the QTM times for these SMMs vary as 0.28, 1.4 and 6.7 ms when X ¼ Cl, Br or I, respectively. The coordination polymer [(Cp )2Dy(m-Cl)2K(THF)]1 shows a surprisingly high barrier of Ueff ¼ 379 cm−1 and relatively long tQTM of 70 ms, with the potassium cations helping to reduce the extent to which the chloride ligands contribute to the equatorial crystal field. Consistent with observations on [Cp’2Dy(m-CH3)]2, slow relaxation of the magnetization in [Cp’2Dy(Z3-C3H5)] can only be observed in an applied DC field of Hdc ¼ 4000 Oe owing to the strong equatorial crystal field arising from the allyl ligand. The importance of a weak equatorial crystal field to complement the strong axial crystal field generated within the dysprosium metallocene building block was further underscored by the SMM properties of [(Cp )2Ln(m-Ph2BPh2)] with Ln ¼ Tb or Dy and their magnetically dilute analogues (Fig. 4). Barriers of 216 and 331 cm−1 were determined for the diluted terbium and dysprosium versions, respectively, in applied DC fields.60 An important result arising from the study of the [(Cp )2Dy(X)(THF)] SMMs was the proposal that an ion-separated species containing the cation [(Cp )2Dy]+ should have excellent SMM properties that approach the axial limit for dysprosium.59 In parallel, a similar conclusion based on a cation of the type [Cp2Dy]+ was drawn following a study of [(Cp )2Dy(m-Fp)]2 (Fp ¼ CpFe(CO)2).61 The synthesis of the isocarbonyl-bridged bimetallic compound was achieved in the salt metathesis reaction between [(Cp )2Dy(m-Ph2BPh2)] and KFp. The resulting SMM displayed an energy barrier of 662(2) cm−1, with the relaxation below 10 K dominated by QTM. The larger barrier is thought to be a consequence of the weak equatorial crystal field provided by the isocarbonyl oxygen donor atoms. Relaxation at higher temperature is then characterized by Raman and Orbach processes and/or thermally assisted QTM (Fig. 6). A magnetic dilution experiment, in which the isostructural mono-dysprosium

O Dy

C

Fe

Energy /cm-1

800 600 │±11/2)

400

│±13/2)

200

│±15/2)

0 –10

5

0

5

10

Magnetic moment /mB Fig. 6 Molecular structure of [(Cp )2Dy(m-Fp)]2 and mechanism for relaxation of the magnetization in individual Dy3+ ions. Red arrows represent transitions between states: more probable transitions are represented by dark shading.

218

f-Element Organometallic Single-Molecule Magnets

compound [(Cp )2Dy(m-Fp)2Y(Cp )2] was dispersed in a matrix of the isostructural diamagnetic compound [(Cp )2Y(m-Fp)]2, revealed no significant differences in the high-temperature relaxation processes, but a lengthening of tQTM from 0.23(2) s to 17(3) s. The dilution study highlighted the detrimental effects of weak magnetic exchange on the SMM properties. The hysteresis loops of the non-dilute and dilute samples were also very similar, with butterfly-shaped loops remaining open up to 6.6 K with an average sweep rate of approximately 20–24 Oe s−1. A theoretical study revealed that the thermally activated relaxation should proceed via at least the fourth-excited KD or possibly a higher-lying state (Fig. 6). The sensitivity of the electronic structure of Dy3+ to even minor changes in the coordination environment was demonstrated through a study of the coordination polymer [Cp 2Dy(m-OC)W(Cp)(CO)(m-CO)]1 (Fig. 4).62 The relaxation in this SMM is characterized by an energy barrier of 557(18) cm−1 and tQTM ¼ 3.7 ms. A computational study of an extended fragment of the polymer considered the tungsten-to-CO back-bonding and how this classical organometallic bonding motif impacts the equatorial crystal field experienced by Dy3+, the main conclusion being that electron correlation effects outside of the lanthanide 4f orbital manifold can be important. The related monometallic dysprosium isocarbonyl complex [{CpW(CO)2(m-CO)}3Dy(THF)5] was synthesized in the deprotonation reaction of [Dy{N(SiMe3)2}3] with [CpW(CO)2H].63 The geometry of the Dy3+ center is a distorted square antiprism, which enables slow relaxation of the magnetization to be observed under an applied DC field of 400 Oe, leading to an energy barrier of 12.6 cm−1. In the series of isostructural pnictogen-bridged trimetallic SMMs [(CpMe)2Dy{m-E(H)Mes}]3 (Fig. 7) (E ¼ P, As, Sb; Mes ¼ mesityl), the metallocene units adopt local C2v symmetry and display remarkably similar bond lengths and angles, but the distances to the bridging phosphide, arsenide- and stibinide ligands vary in the ranges 2.920(6)–2.946(6) A˚ , 2.984(2)–3.012(2) A˚ and 3.118(2)–3.195(2) A˚ , respectively.64–66 In the case of the phosphide-bridged metallocene [(CpMe)2Dy{m-P(H)Mes}]3, maxima in the w”(n) data were observed in the temperature range 2–31 K, resulting in Ueff ¼ 210(6) cm−1 with t0 ¼ 6.53  10−9 s. The temperature dependence of ln t on T−1 reflects that the relaxation is dominated by Orbach processes, with minimal contributions from QTM even at low temperatures. A theoretical study suggested that the dominant thermal relaxation process in this SMM occurs via the second-excited KD. An increase in the barrier to 256(6) cm−1 was observed in a magnetically dilute sample. Similar analysis of the arsenide-bridged compound [(CpMe)2Dy{m-As(H)Mes}]3 yielded an energy barrier of 256(5) cm−1 for the non-dilute compound and 301(9) cm−1 for the magnetically dilute version, with t0 ¼ 2.01  10−9 s and 4.77  10−7 s, respectively. The SMM properties of the isostructural selenolate-bridged SMM [(CpMe)2Dy(m-SeMes)]3 were measured as Ueff ¼ 285(4) cm−1 with t0 ¼ 2.50  10−8 s, changing to Ueff ¼ 301(7) cm−1 with t0 ¼ 4.48  10−10 s upon dilution.65 A barrier of 345 cm−1 was extracted for non-dilute and dilute versions of [(CpMe)2Dy{m-Sb(H)Mes}]3, with the other fitting parameters showing only minor variations between the two materials. Orbach relaxation in the arsenic- and antimony-bridged SMMs was also calculated to proceed via the second-excited KD. The gradual increase in the barrier on replacing phosphorus by the heavier congeners in group 15 reflects a diminishing influence of the equatorial pnictogen ligands, consistent with the general picture for SMM behavior in dysprosium metallocenes. A theoretical study of the magnetic exchange in [(CpMe)2Dy{m-P(H)Mes}]3 revealed weak antiferromagnetic interactions between the Dy3+ ions, resulting in a six-fold degenerate set of frustrated exchange KDs, depicted in Fig. 8. Reversal or flipping of the magnetic dipoles on the individual dysprosium centers in this compound should, therefore, be more rapid in the non-dilute than in the dilute version, which has a {DyY2} core, meaning that the exchange is removed. These phenomena have important consequences for the magnetic hysteresis properties of the compounds: in the case of the non-dilute phosphide-bridged compound, very narrow M(H) loops were observed at 1.8 K, but in the dilute version, butterfly-shaped hysteresis loops were observed up to 4.4 K with a sweep rate of 26 Oe s−1 (Fig. 8). The same effect of dilution on the magnetic hysteresis was also seen with [(CpMe)2Dy {m-As(H)Mes}]3, [(CpMe)2Dy(m-SeMes)]3 and [(CpMe)2Dy{m-As(H)Mes}]3, with the dilute versions producing open M(H) loops up to 5.4, 5.4 and 4.7 K, respectively, with sweep rates of 31.4, 37.5 and 28 Oe s−1.

Fig. 7 Structures of trimetallic dysprosium metallocene SMMs.

f-Element Organometallic Single-Molecule Magnets

3+

1–

2–

1+

Magnetisation /mB

0.8

219

2-Y2Dy

0.4

0.0

1.8 K 2.2 K 2.6 K 3.2 K 3.6 K 4.0 K 4.4 K

1.8 K

–0.4

2+ 3–

–0.8 –10,000

–3,000 –1,500

–5,000

0

0

5,000

1,500

3,000

10,000

Field / Oe Fig. 8 Left: schematic illustration of the three low-lying, quasi-degenerate exchange levels in [(CpMe)2Dy{m-P(H)Mes}]3 (Dy ¼ blue, P ¼ brown). Right: Magnetic hysteresis loops for the magnetically dilute version of the same compound at the temperatures indicated, using an average scan rate of 26 Oe s−1. Adapted with permission from reference Pugh, T.; Tuna, F.; Ungur, L.; Collison, D.; McInnes, E. J. L.; Chibotaru, L. F.; Layfield, R. A. Nat. Commun. 2015, 6, 7492.

The addition of three equivalents of nBuLi to [(CpMe)2Dy{m-E(H)Mes}]3 (E ¼ P or As) produced the phosphinidene- and arsinidene-bridged complexes [{(CpMe)2Dy(m-EMes)}3Li]2− as salts of two [Li(THF)4]+ cations (Fig. 7). The core geometries of these compounds are qualitatively similar to those of their precursor compounds, with the additional lithium cation residing above the center of the {Dy3E3} rings and bound to the pnictogen donor atoms. The dianionic charge on the phosphinidene and arsinidene ligands results in Dy–E distances of 2.7850(15)-2.8249(15) A˚ and 2.8515(6)-2.8908(7) A˚ , which are much shorter than those in the respective precursor compounds. Concomitantly, lengthening of the DydC distances to the CpMe ligands by approximately 0.05 A˚ also occurs. The shorter distances to the bridging equatorial ligands equate to a strengthening of the equatorial crystal field and the longer distances to the axial CpMe ligands diminish the axiality of the crystal field. These structural changes are expected to result in much poorer SMM properties, which was indeed observed for both complexes. The maxima in w”(n) were observed in the temperature range 1.8–4.2 K and 1.8–5 K for the phosphinidene- and arsinidene-bridged SMMs, respectively, and the barriers were determined to be 13(1) cm−1 and 23(2) cm−1 with t0 ¼ 7.75  10−7 s and 2.99  10−7 s, respectively. A slight increase in the energy barrier for [Li(THF)4]2[{(CpMe)2Dy(m-AsMes)}3Li] to 35(2) cm−1 was observed, although a barrier for the magnetically dilute phosphinidene-bridged SMM could not be determined. Attempts to synthesize a stibinidene-bridged trimetallic SMM from [(CpMe)2Dy{m-Sb(H)Mes}]3 instead produced the unusual compound [{(CpMe)2Dy{m-(SbMes)3Sb}], in which the three Dy3+ ions are bridged by a Zintl-like [Sb4Mes3]3− anion.66 The same compound can also be synthesized by adding one equivalent of MesSbH2 to [(CpMe)2Dy{m-E(H)Mes}]3 in a reaction that proceeds via a dehydrocoupling mechanism, accompanied by formation of H2 and mesitylene (Scheme 1). A decrease in the energy barrier to Ueff ¼ 272 cm−1 occurs, with only slight variations in the SMM properties upon measuring the magnetically dilute version. The lower barrier for [{(CpMe)2Dy{m-(SbMes)3Sb}] is thought to be a consequence of the slightly shorter, on average, DydSb distances and slightly greater calculated charges on the antimony atoms, consistent with the general magneto-structural correlation developed to explain the SMM properties of dysprosium metallocenes.

Scheme 1 Synthesis of [{(CpMe)2Dy{m-(SbMes)3Sb}].

The reactions of the pyridyl-bridged ansa-metallocene salt di-sodium 2,6-bis(cyclopentadienylmethyl)pyridine (Napy 2 Cp2) with Dy(OTf )3 or DyCl3 produced the triflate- and chloride-bridged dimers [(pyCp2)Dy(m-X)]2, where X ¼ TfO− or Cl− and the triflate ligands coordinate using one oxygen donor per dysprosium (Fig. 9).67 The subsequent reaction of [(pyCp2)Dy(m-OTf )]2 with THF leads to formation of the monometallic complex [(pyCp2)Dy(m-OTf )(THF)]. In each of the three complexes, the ansa-metallocene ligand binds through two Z5-Cp ligands and the N-donor of the pyridyl group, resulting in coordination numbers of nine when considering the bridging or THF ligands. An important structural difference between [(pyCp2)Dy(m-OTf )]2 and [(pyCp2)Dy(m-Cl)]2 is the large difference in the Dy


Dy distances of 6.068(1) A˚ and 4.252(1) A˚ , respectively.

220

f-Element Organometallic Single-Molecule Magnets

Fig. 9 Molecular structures of [(pyCp2)Dy(m-OTf )]2 (left) and [(pyCp2)Dy(m-OTf )(THF)] (right) (grey ¼ C, blue ¼ N, red ¼ O, bright green ¼ F, yellow ¼ S, dark green ¼ Dy). The solid red lines indicate the easy-axis of magnetization in the ground KD. Reproduced with permission from reference Burns, C. P.; Wilkins, B. O.; Dickie, C. M.; Latendresse, T. P.; Vernier, L.; Vignesh, K. R.; Bhuvanesh, N. S.; Nippe, M. Chem. Commun. 2017, 53, 8419–8422.

For both triflate-ligated complexes, the relaxation is dominated by QTM at low temperatures (below 4–6 K) before thermal relaxation via Orbach processes is observed at higher temperatures. Identical energy barriers of Ueff ¼ 49 cm−1 were determined for the two complexes, along with t0 ¼ 4.8  10−7 s and 7.2  10−7 s for the dimer and monomer, respectively. In contrast, it was not possible to determine an energy barrier for the chloride bridged dimer in zero DC field or in applied fields. A theoretical study revealed the exchange coupling in both dimers to be very weak, such that the SMM properties arise from single-ion effects. In all cases, the dominant relaxation process at higher temperatures was determined to proceed via thermally assisted QTM involving the first-excited KD. The reaction of [(pyCp2)Dy(m-OTf )(THF)] with K[MCp(CO)2] (M ¼ Fe, Ru) produced [(pyCp2)Dy–MCp(CO)2], which contain unsupported metal-metal bonds, with DydFe and DydRu distances of 2.884(2) A˚ and 2.9508(5) A˚ , respectively.68 Theoretical studies of the metal-metal bonds indicated strong s-donation from the transition metal to the lanthanide, a finding supported by the 57Fe Mössbauer spectrum of the iron-containing version. Donation from iron to dysprosium was calculated to be stronger than with ruthenium, according to results from QTAIM calculations (quantum theory of atoms in molecules). Slow relaxation of the magnetization was not observed for either compound in the absence of an applied DC field. However, in applied fields of 1500 Oe and 1600 Oe, respectively, maxima in w”(n) were observed in the range 3–10 K, which led to energy barriers of 43 cm−1 and 46 cm−1 being determined. These results indicated that the crystal field splitting in both compounds is similar. The suitability of complexes of the type [Cp2Dy(m-X)]n for SMM behaviour furnishes a consistent picture of the slow magnetic relaxation being a consequence of the anisotropic oblate 4f electron density being sandwiched between the cyclopentadienyl ligands, which provide a strong axial crystal field. As the influence of the equatorial ligands grow, the SMM properties tend to diminish. In a complementary manner, support for this axial theory was provided by an SMM containing the complex anion [Cp2Er(OtBu)2]−. The prolate electron density of Er3+ is not expected to respond well to the axial crystal field, and this was indeed observed with small energy barrier of 20 cm−1 in an applied field of 1500 Oe.69 The basic lanthanide metallocene blueprint has been used to develop structurally similar sandwich complexes containing ligands with differing properties to those offered by the familiar cyclopentadienyl ligand. For example, the open pentagonal face of the carborane dianion nido-[C2B9H11]2− allows Z5-coordination to Dy3+, as found in the sandwich complex [(Z5-C2B9H11)2Dy (THF)2]− (Fig. 10).70 The w”(n) data for this complex displayed maxima up to 40 K, with the high-temperature relaxation indicating Orbach type relaxation, with the lower-temperature relaxation thought to consist of Raman and QTM processes. An energy barrier of 299(3.5) cm−1 was determined with t0 ¼ 1.2(3)  10−9 s. Waist-restricted hysteresis loops were observed up to 4.9 K (sweep rate 15 Oe s−1) for the magnetically dilute version. Theoretical analysis of this SMM revealed that the easy axis of magnetization is oriented towards the carborane ligands. The three lowest-lying KDs are highly axial and the most probable thermal relaxation process proceeds via the fourth and/or fifth KDs. Using the hypothetical model complex [(Z5-C2B9H11)2Dy]− with the centroidDy-centroid angle constrained to be 180 , an energy barrier in the region 1520–1764 cm−1 was predicted. An attempt to block coordination of THF from the coordination sphere of dysprosium by using a bulkier, substituted derivative of the carborane led to formation of a half-sandwich complex with Z5-coordination by one [C2B9H11]2− ligand in addition to coordination by two more carborane ligands via bridging hydrides. The effect of replacing one Z5-carborane ligand with two m-hydride carborane ligands is to increase the energy barrier to 559(5) cm−1. The theoretical analysis predicted thermal relaxation via a bunched group of higher-lying KDs, possibly up to the seventh KD. Butterfly-shaped magnetic hysteresis loops were observed up to 6.8 K in the magnetically dilute analogue of this SMM using an average sweep rate of 15 Oe s−1. The eight-membered ring pentalene (Pn) can be regarded as two cyclopentadienyl rings fused along a carbon-carbon bond. The pentalene dianion is a 10p aromatic system, with several substituted versions being able to form Z8-complexes with lanthanides and actinides.71 The bulky pentalene ligand [1,4-(iPr3Si)2C8H4]2− (Pn{) in the sandwich complex [(Z8-Pn{)Dy(Z5-Cp )] coordinates via a geometry in which the two five-membered rings fold towards the metal (Fig. 11).72 Electronic structure calculations showed that the ground KD is highly axial in character and that the easy-axis of magnetization in this KD passes through the center of the Cp ligand and the midpoint of the pentalene ligand (the so-called ‘body’ carbon atoms). The SMM properties consist of an energy barrier of 188(11) cm−1, which increases to 245(28) cm−1 upon dilution. The dominant thermal relaxation in this system proceeds

f-Element Organometallic Single-Molecule Magnets

221

Fig. 10 Structure of the [(Z5-C2B9H11)2Dy(THF)2]− anion with thermal ellipsoids at the 30% probability level. Reproduced with permission from reference Jin, P.-B.; Zhai, Y.-Q.; Yu, K.-X.; Winpenny, R. E. P.; Zheng, Y.-Z. Angew. Chem. Int. Ed. 2020, 59, 9350–9354.

Fig. 11 Left: molecular structure of [(Z8-Pn{)Dy(Z5-Cp )] and frequency-dependence of the imaginary component of the AC susceptibility at temperatures in the range 2–41 K. Reproduced with permission from reference Kilpatrick, A. F. R.; Guo, F.-S.; Day, B. M.; Mansikkamäki, A.; Layfield, R. A.; Cloke, F. G. N. Chem. Commun. 2018, 54, 7085–7088.

via the first-excited KD. The pentalene ligand provides the crystal field with an axial component via the two body carbon atoms and, simultaneously, a non-negligible equatorial component via the wing-tip carbon atoms, reminiscent of Z8-cyclooctatetraenyl ligands (see Section 1.08.2.2), making it poorly suited to the design of SMMs with enhanced properties.

1.08.2.1.2

Lanthanide half-sandwich complexes as SMMs

The heterobimetallic cage complex [(Cp Dy)6Cl16K4(THF)6] forms in the reaction of DyCl3 with KCp in a 1:1 stoichiometric ratio (Fig. 12).73 The elaborate structure of this molecule consists of six structurally similar {Cp DyCl5} units bridged by the chloride ligands. A theoretical study revealed that the easy-axis of magnetization in the ground KD (which is assumed to have dominant |MJ | ¼ 15/2 character in this model) is oriented towards the Cp ligands. The w”(T ) data in zero DC field displays well-defined maxima in the frequency range n ¼ 1–1488 Hz, although with sharp rises in w” at low temperatures for the higher-frequency data, indicating the onset of QTM. In an applied DC field of 1500 Oe, the maxima in the imaginary component of the AC susceptibility move to slightly higher temperatures and the QTM is effectively suppressed. In zero-field, fitting the relaxation time data with Orbach, Raman and QTM terms allowed an energy barrier of 316 cm−1 to be determined, with t0 ¼ 1.7  10−11 s and tQTM ¼ 2.1 s. The in-field data were fitted without a QTM term, which yielded Ueff ¼ 390 cm−1 and t0 ¼ 1.3  10−12 s. The M(H) hysteresis for this polymetallic complex also displayed waist-restricted hysteresis loops up to 4.5 K when using a sweep rate of 20 Oe s−1. The dysprosium half-sandwich complexes [(CpR)Dy(DBM)2(THF)], where CpR is Cp , Cn5Pr4Ph or C5Me4SiMe3 and DBM is the bidentate b-diketonate ligand dibenzoylmethanoate, consist of formally 8-coordinate Dy3+ ions, with a two-fold rotation axis being

222

f-Element Organometallic Single-Molecule Magnets

Fig. 12 Molecular structure of [(Cp Dy)6Cl16K4(THF)6] and w”(T ) data collected in an applied field of 1500 Oe. Reproduced with permission from reference Wu, J.; Demeshko, S.; Dechert, S.; Meyer, F. Chem. Commun. 2020, 56, 3887–3890.

Fig. 13 Molecular structure of [(Cp )Dy(DBM)2(THF)] showing the easy axis of magnetization calculated by two different models, i.e. ab initio (red line) and electrostatic (blue line). Reproduced with permission from reference Long, J.; Shestakov, B.G.; Liu, D.; Chibotaru, L.F.; Guari, Y.; Cherkasov, A.V.; Fukin, G.K.; Trifonov, A.A.; and Larionova, J. Chem. Commun. 2017, 53, 4706–4709.

approximately coincident with the OTHF-Dy-CpR axis (Fig. 13).74 The three DBM-ligated complexes display SMM behavior in zero DC field, with the properties of [(Cp )Dy(DBM)2(THF)] and [(Cn5Pr4Ph)Dy(DBM)2(THF)] being qualitatively similar. For the former complex, the w”(n) data show temperature-independent maxima below 6 K and only two maxima above 10 K; for the latter complex, a similar frequency independence of the susceptibility is observed below 8 K, but with more w”(n) maxima above 10 K. The behavior of these two complexes reflects QTM at low temperatures followed by Orbach relaxation at higher temperatures, a hypothesis confirmed with suppression of the QTM in an applied field of 2000 Oe. In contrast, the w”(n) data for [(C5Me4SiMe3)Dy(DBM)2(THF)] displays temperature-dependent maxima at all temperatures in the range 9–22 K. The contrasting behavior of [(C5Me4SiMe3)Dy(DBM)2(THF)] indicates dominant relaxation via activated processes and also that QTM is not significant in this system. Analysis of the relaxation time data for [(Cp )Dy(DBM)2(THF)] and [(Cn5Pr4Ph)Dy(DBM)2(THF)] produced energy barriers of Ueff ¼ 32 cm−1 and 53 cm−1, respectively, with t0 ¼ 7.8  10−6 s and t0 ¼ 2.6  10−6 s. The curved nature of the T−1 dependence of ln t for [(C5Me4SiMe3)Dy(DBM)2(THF)] suggested a combination of Raman and Orbach processes, and analysis of the relaxation times using both terms led to an energy barrier of 222 cm−1 with t0 ¼ 6.3  10−11 s. Consistent with the SMM parameters determined from AC susceptibility measurements, the M(H) hysteresis loops measured at 10 Oe s−1 are closed at 1.8 K for [(Cp )Dy(DBM)2(THF)], slightly open but waist-restricted at 1.8 K for [(Cn5Pr4Ph)Dy (DBM)2(THF)], and waist-restricted up to 4 K for [(C5Me4SiMe3)Dy(DBM)2(THF)]. Theoretical analysis of all three complexes revealed a complex picture in which the substituents on the cyclopentadienyl ligands play an important role in determining the SMM properties. The dipotassium salt of the diazadiene (DAD) ligand 2,6-iPr2C6H3NC(H)]C(H)C6Hi3Pr2–2,6 reacts with DyCl3 followed by the addition of MCp (M ¼ Li or K) to give the half-sandwich complexes [Cp Dy(DAD)L] in which L ¼ THF when M ¼ K, and L ¼ ClLi(THF)3 when M ¼ Li.75 The six-coordinate dysprosium ions occupy piano-stool type geometries, leading to broad curves with poorly defined maxima in the w”(n) data in zero DC field below 20 K. Clearer maxima became discernible in applied fields of 2000 Oe and 1000 Oe for [Cp Dy(DAD)(THF)] and [Cp Dy(DAD){ClLi(THF)3}], respectively, allowing barriers of 206  10 cm−1 and 20  4 cm−1 to be determined, with t0 ¼ (10  4)  10−9 s and t0 ¼ (5  3)  10−9 s. The Cp and DAD ligand both strongly influence the orientation of the easy axis of magnetization in these complexes, with the additional THF and {ClLi(THF)3} ligands providing competing equatorial crystal fields.

f-Element Organometallic Single-Molecule Magnets

223

Fig. 14 Molecular structure of [Dy(Z5-tBu4Carb)(o-Me2NC6H4CH2)2]. Reproduced with permission from reference Long, J.; Selikhov, K. A.; Lyssenko, K. A.; Guari, Y.; Larionova, J.; Trifonov, A. A. Organometallics 2020, 39, 2785–2790.

Analogues of half-sandwich complexes containing N-donor or pyrrolyl-type ligands are uncommon.76,77 With appropriate steric bulk, pyrazolyl ligands can be forced to adopt an Z5-bonding mode to dysprosium, as illustrated with the half-sandwich complex [Dy(Z5-tBu4Carb)(o-Me2NC6H4CH2)2], in which tBu4Carb is 1,3,6,8-tetra-tert-butyl-9H-carbazole and o-Me2NC6H4CH2 is N,N-dimethylaminobenzyl (Fig. 14).78 The piano-stool complex contains a 7-coordinate dysprosium center with the Dy3+ cation residing above the center of the pyrrole ring. The w”(n) data, which show a series of maxima in the temperature range 2–22 K in zero DC field, were interpreted with a model invoking only Raman and QTM relaxation processes.

1.08.2.1.3

Lanthanide metallocene SMMs with radical bridging ligands

A notable trend among the family of [Cp2Dy(m-X)]n SMMs is that, while large effective energy barriers are possible, the magnetic hysteresis properties typically consist of S-shaped or ‘waist-restricted’ loops with negligible coercivity and remanent magnetization. A step-change in the understanding of hysteresis in SMMs occurred in 2011 with reports of the di-lanthanide compounds [K(18crown-6)(THF)2][Ln2{N(SiMe3)2}4(THF)2(m:Z2:Z2-N2)], where the lanthanide is either terbium or dysprosium and the S ¼ ½ radical trianion [N2]3− bridges between the Ln3+ ions (Fig. 15).79–81 The main difference in the M(H) hysteresis properties was the observation of open loops with appreciable coercive fields that remain open up to 14 K and 8.3 K for the terbium and dysprosium versions, respectively. The terbium radical-bridged complex displayed a coercive field of a remarkable 5 T at 11 K when using an average sweep rate of 9 Oe s−1 and the dysprosium congener has a coercive field of 1.5 T up to 6 K with a sweep rate of 800 Oe s−1. An indication of the strength of the anisotropic exchange between the metal ions and the radical ligand in the radical-bridged terbium- and dysprosiumamido complexes was obtained from the isotropic exchange coupling constant, J, in the isostructural gadolinium compound [K(18crown-6)(THF)2][Gd2{N(SiMe3)2}4(THF)2(m:Z2:Z2-N2)]. Using a spin Hamiltonian where the exchange parameter is expressed as a –2J term, the coupling constant was determined to be −27 cm−1, i.e. considerably stronger than a typical exchange coupling constant in a multimetallic lanthanide complex (i.e. on the order of 0.1–1 cm−1).4 The origins of the radical-bridge effect can be assigned to the unpaired spin on the ligand effectively acting in lieu of an external magnetic field, which lifts the degeneracy of the two components of each Kramers doublet, therefore reducing the probability of QTM. However, despite the hysteresis properties of these SMMs, the effective energy barriers are modest, being determined as Ueff ¼ 227.0(4) cm−1 with t0 ¼ 8.2(1)  10−9 s and 123 cm−1 with

Fig. 15 Molecular structure of [Tb2{N(SiMe3)2}4(THF)2(m:Z2:Z2-N2)]− and magnetic hysteresis properties at the temperature indicated using a scan rate of 9 Oe s−1 (C ¼ grey, N ¼ blue, O ¼ red, Si ¼ green, Tb ¼ brown). Reproduced with permission from reference Rinehart, J.D.; Fang, M.; Evans, W.J.; Long, J.R. J. Am. Chem. Soc. 2011, 133, 14236–14239.

224

f-Element Organometallic Single-Molecule Magnets

Fig. 16 Left: molecular structure of [(Cp 2Dy)2(m-bpym)]+ indicating the exchange interactions. Right: Magnetic hysteresis loops at the temperatures indicated with a scan rate of 30 Oe s−1 and (inset) temperature dependence of the relaxation time. Reproduced with permission from reference Demir, S.; Zadrozny, J. M.; Nippe, M.; Long, J. R. J. Am. Chem. Soc. 2012, 134, 18546–18549.

t0 ¼ 8  10−1 s, respectively. The energy barriers are most likely to be a consequence of the relatively low symmetry crystal fields in addition to the fact that the [N2]3− ligand also provides a competing equatorial crystal field. Further exploration of the use of radical bridging ligands in lanthanide metallocene SMMs led to the characterization of [(Cp 2Ln)2(m-bpym)][BPh4] (Ln ¼ Gd, Tb, Dy), where the Ln3+ ions are bridged by the S ¼ ½ radical anion of bipyrimidyl (bpym) (Fig. 16).82 Insight into the metal-ligand exchange interaction was again obtained from magnetic susceptibility measurements on the gadolinium version, which yielded a ground state spin of Stot ¼ 13/2 and an isotropic exchange coupling constant of J ¼ −10 cm−1 (−2J formalism). The coercive field for the dysprosium-bpym complex at 3 K was measured as 0.6 T with a scan rate of 20 Oe s−1, with the loops remaining open up to 6.5 K. As with the radical-bridged amido complexes, maxima in the w”(n) data were observed in the range 7.5–17.5 K in zero DC field, and the effective energy barrier was again modest at only 87.8(3) cm−1 with t0 ¼ 1.03(4)  10−7 s. Analogous characterization of the bpym-bridged terbium complex produced a barrier of 44(2) cm−1 with t0 ¼ 4(1)  10−8 s. A recent study on a series of isostructural bipyrimdyl-bridged dilanthanide SMMs has also shown that the SMM properties can be influenced by the electronic properties of the substituents on the bpym ligand.83 In the trimetallic radical-bridged compound [(Cp 2Dy)3(HAN)], where HAN is the tritopic phenanthroline-like ligand hexaazatrinaphthylene, an effective energy barrier of 51 cm−1 with t0 ¼ 1.2  10−8 s was determined along with open magnetic hysteresis loops up to 3.5 K.84 Refluxing the di-lanthanide compounds [K(2.2.2-crypt)(THF)][{(C5Me4H)2Ln(THF)}2(m-N2)] in 2-methyl-THF resulted in removal of the THF ligands to give [K(2.2.2-crypt)][{(C5Me4H)2Ln}2(m-N2)] (Ln ¼ Gd, Tb, Dy), in which the lanthanide centres occupy more symmetrical environments.85 In both sets of complexes, the lanthanides are bridged by radical [N2]3− ligands. The structural changes occurring upon desolvation of the Ln3+ ions are thought to facilitate more effective exchange interactions aross the {Ln(m-N2)Ln} units. For example, in [{(C5Me4H)2Tb(THF)}2(m-N2)]− the {Tb(m-N2)Tb} dihedral angle is 173.45(16) , whereas in [{(C5Me4H)2Tb}2(m-N2)]− the analogous angle is 178.5(2) . The effective energy barriers for the di-terbium complexes were determined to be 242(2) cm−1 with t0 ¼ 4(1)  10−8 s for the THF-solvate and 276(1) cm−1 with t0 ¼ 1.3(1)  10−7 s for the THF-free complex. A second Orbach process was also identified for the THF-free complex with a barrier double that of the first Orbach process, i.e. Ueff ¼ 564(17) cm−1 with t0 ¼ 2(1)  10−11 s, which could be explained by a spin ladder structure to the excited exchange states. The hysteresis properties of the THF-solvated di-terbium complex include open M(H) loops up to 15 K with a coercive field of 3.7 T at 11 K (scan rate of 100 Oe s−1). In contrast, a very large coercive field of 7.9 T was measured for the THF-free diterbium SMM at 10 K, with the loops remaining open down to 2 K. The giant coercive field measured in the di-terbium SMM represents, at the time of writing, the strongest yet observed in an SMM. In contrast, the isostructural dysprosium complex [{(C5Me4H)2Dy(THF)}2(m-N2)]− does not produce open hysteresis loops, whereas [{(C5Me4H)2Dy}2(m-N2)]− displays hysteresis up to 8 K, with a coercive field of 1 T at 5.5 K. The effective energy barriers for the di-dysprosium complexes are also markedly lower than those of the terbium analogues, being 110(1) cm−1 with t0 ¼ 3.1(1)  10−9 s and 108.1(2) cm−1 t0 ¼ 1.7(1)  10−8 s for the THF-solvated and unsolvated versions, respectively. Although careful use of radical ligands as bridges between lanthanide ions can be a highly effective strategy for improving the hysteresis properties, the symmetry of the individual lanthanide coordination environments and the balance between axial and equatorial contributions to the crystal fields are also important considerations. This point was illustrated in the series of indigo-bridged dimetallic complexes [{Cp 2Ln}2(m-ind)]n– (n ¼ 0, 1, 2; Ln ¼ Gd, Dy), where the indigo ligand can be present in the diamagnetic dianionic form, the trianionic S ¼ ½ form, or the diamagnetic tetranionic form.86 The structures of these dimetallic compounds consist of two Z5-Cp ligands, with the indigo ligand binding via the unsymmetrical N,O-mode to each metal center. The redox non-innocent nature of indigo allows consecutive one- and two-electron reductions via addition of KC8 to [{Cp 2Ln}2(m-ind)], resulting in the formation of [{Cp 2Ln}2(m-ind)]− and [{Cp 2Ln}2(m-ind)]2− with solvated potassium counter-cations (Scheme 2).

f-Element Organometallic Single-Molecule Magnets

Ln

2

-2 C3H6

LnCp*2

O

H2ind

N 2 N O

Cp*2Ln 2 KC8 / thf

O

(thf)3K

225

KC8 / thf

LnCp*2

O

LnCp*2 N

N 3

4

N

N Cp*2Ln

O

K(thf)3

Cp*2Ln

O

[K(thf)6] Scheme 2 Synthesis of indigo-bridged di-lanthanide compounds. Reproduced with permission from reference Guo, F.-S.; Layfield, R. A. Chem. Commun. 2017, 53, 3130–3133.

The dimetallic gadolinium complex [{Cp 2Gd}2(m-ind)]− contains Gd3+ ions directly coupled to the S ¼ ½ radical indigo ligand with J ¼ −11 cm−1 (−2J formalism), whereas only weak antiferromagnetic exchange with J ¼ −0.013(1) cm−1 and − 0.018(1) cm−1 between the metal ions was observed in [{Cp 2Gd}2(m-ind)] and [{Cp 2Gd}2(m-ind)]2−, respectively. Of the three indigo-bridged dysprosium complexes, [{Cp 2Dy}2(m-ind)] and [{Cp 2Dy}2(m-ind)]− displayed SMM behavior whereas [{Cp 2Dy}2(m-ind)]2− did not. In the case of [{Cp 2Dy}2(m-ind)], maxima in w”(n) were observed in the temperature range 5–18 K, leading to an energy barrier for the Orbach process of 39(1) cm−1 with t0 ¼ 5.08  10−5 s. In [{Cp 2Ln}2(m-ind)]−, the maxima were observed in the region 2.7–5.4 K and Ueff was determined to be 35(1) cm−1 with t0 ¼ 1.60  10−8 s. Both SMMs also displayed narrow, S-shaped hysteresis loops at 1.8 K with negligible coercivity. The relatively poor SMM properties are fully consistent with the magneto-structural correlation developed previously for multimetallic dysprosium metallocene SMMs, whereby the hard donor atoms of the bridging indigo ligands provide a strong, competing equatorial crystal field that limits the effective barrier height. As the charge on the indigo ligand increases across the series, the equatorial contribution strengthens, hence the SMM properties diminish. Consequently, it was proposed that unpaired spin density on the bridging ligands does not necessarily lead to improved SMM properties: the symmetry and strength of the crystal field are also important. The bis(terdentate) ligand 2,3,5,6-tetra(2-pyridyl)pyrazine (tppz) bridges between two lanthanide metallocene units to give the dimetallic complexes [{Cp 2Ln}2(m-tppz)]+ as salts of [BPh4]−, where Ln ¼ Gd, Tb and Dy and the heterocyclic ligand is present as the radical monoanion with S ¼ 1/2.87 Subsequent reduction of the di-lanthanide cations with two equivalents of KC8 produced [K(2.2.2-crypt)][{Cp 2Ln}2(m-tppz)], in which the trianionic [tppz]3− ligand is also in the S ¼ ½ form (Fig. 17). Regarding the [Cp ]− as occupying three coordination sites, each lanthanide ion in the dimetallic complexes is formally 9-coordinate. The digadolinium complexes have Stot ¼ 13/2 magnetic ground states, with exchange coupling constants of J ¼ −6.91(4) cm−1 and − 6.29 (3) cm−1 (−2J formalism) for the cation and anion, respectively. Evidence for stronger coupling in the terbium and dysprosium versions based on the DC magnetic susceptibility measurements was proposed. The similar values of J for the di-gadolinium

Fig. 17 Molecular structure of [{Cp 2Ln}2(m-tppz)]+ and the w00 (n) data in zero applied field. Reproduced from reference Demir, S.; Nippe, M.; Gonzalez, M. I.; Long, J. R. Chem. Sci. 2014, 5, 4701–4711.

226

f-Element Organometallic Single-Molecule Magnets

compounds was interpreted with the aid of DFT calculations, which revealed similar amounts of unpaired spin density on the nitrogen donor atoms. The greater formal charge on the [tppz]3− ligand in the anionic terbium and dysprosium complexes provides an effective competing equatorial crystal field, resulting in a lack of slow relaxation of the magnetization. In contrast, the lower charge on the [tppz]− ligand reduces the impact of the equatorial crystal field, allowing maxima in w”(n) to be observed in zero DC field from 1.8–2.45 K for the di-terbium complex cation and from 3.5–8 K for the complex anion. The magnetic relaxation time for both cations displays a linear dependence on T−1, indicating dominant relaxation via Orbach processes and minimal QTM. A very small barrier of 5.1(1) cm−1 with t0 ¼ 6(1)  10−6 s was determined for the terbium cation and 35.9(2) cm−1 with t0 ¼ 2.1(1)  10−7 s for dysprosium. S-shaped hysteresis loops were also observed for the dysprosium version up 3.50 K (sweep rate of 30 Oe s−1). The important role of site symmetry in addition to the presence of a radical ligand was identified as a key part of the design process. The relatively diffuse nature of the unpaired spin density on the radical ligand is also likely to be a contributing factor to the lower blocking temperatures of these SMMs compared to those with bridging [N2]3− ligands.

1.08.2.1.4

Cationic dysprosium metallocene SMMs [(CpR)2Dy]+

The magneto-structural correlation developed for dysprosium metallocene SMMs with the general formula [(CpR)2Dy(m-X)]n paints a clear picture of the effects of the equatorial X ligands, which tend to attenuate the effective energy barrier Ueff and have a negative impact on the magnetic blocking temperature. The proposal in 2016 that complete removal of these equatorial ligands – to give an ion-separated complex with composition [(CpR)2Dy]+59,61 has had a transformative effect on single-molecule magnetism in the form of SMM behavior being observed at unprecedentedly high temperatures. In a qualitative theoretical model, discrete metallocene cations should consist of a purely axial crystal field, provided the counter anion does not coordinate directly to dysprosium. At the time of writing, the SMM that currently represents the leading system in terms of energy barrier and magnetic blocking temperature is the ion-separated compound [(Z5-Cp )Dy(Z5-Ci5Pr5)][B(C6F5)4], obtained using the reaction sequence depicted in Scheme 3.45

Scheme 3 Synthesis of [(Z5-Cp )Dy(Z5-Ci5Pr5)][B(C6F5)4].

The important structural parameters in [(Z5-Cp )Dy(Z5-Ci5Pr5)]+ (the so-called 5 cation) are Dy–Cpcent distances to the Cp and Ci5Pr5 ligands of 2.296(1) A˚ and 2.284(1) A˚ , respectively, with the Cp-Dy-Cp angle being 162.507(1) . As shown below through comparisons with other dysprosocenium cations, the geometric parameters for this cation result in what is, to date, the strongest and the most axial (i.e. closest to 180 ) crystal field yet known in a trivalent dysprosium SMM. The w”(n) data for the 5 cation feature well-defined maxima up to a temperature of 138 K (Fig. 18). The magnetization decay observed for this SMM at 77 K occurs over a period of 50 s, increasing to 500 min at 15 K, and approaching a decay time on the order of 24 h at 2 K. The relaxation time t has a strong linear dependence on T−1 in the region 55–138 K, before moving into a curved dependence at temperatures in the range 10–55 K and then into a temperature independent region at lower temperatures. Fitting of the relaxation time data yielded parameters of Ueff ¼ 1541(11) cm−1, t0 ¼ 4.2(6)  10−12 s and tQTM ¼ 2.5(2)  10−4 s, with the energy barrier representing the current record for all SMMs. A second notable feature of the SMM properties of the 5 cation is that the magnetic hysteresis loops remain open up to 80 K when using a relatively slow scan rate of 25 Oe s−1, making it the first SMM with a blocking temperature to exceed the boiling point

0.02

0.00 0.1

5 4

W/S

0.04

10000 1000 100 10 1 0.1 0.01 1E-3 1E-4 1E-5

M / Nb

Dy1

x´´ M /cm3 mol–1

0.06

1

10

v / Hz

100

1000

3 2 1

T = 2 ~ 75 K 200 Oe/s

0 –1 –2 –3 –4 –5 –60

–40

–20

0

20

40

60

H / kOe

0.0

0.1

0.2

0.3

0.4

0.5

T –1/K–1

Fig. 18 Left: molecular structure of the 5 cation. Centre: the corresponding w”(n) data from T ¼ 82 K (green curve) up to 138 K (blue curve). Right the t(T−1) data and (inset) the magnetic hysteresis data under the conditions indicated. Reproduced with permission from reference Guo, F.-S.; Day, B. M.; Chen, Y.-C.; Tong, M.-L.; Mansikkamäki, A.; Layfield, R. A. Science 2018, 362, 1400–1403.

f-Element Organometallic Single-Molecule Magnets

227

of liquid nitrogen. A study of this SMM using ab initio calculations revealed that the easy-axis of magnetization is roughly oriented towards the centers of the two cyclopentadienyl ligands in the first seven KDs, with a maximum deviation from the ground KD of only 5.3 , confirming the exceptional axiality of this system. Furthermore, each of the six lowest KDs is described by a definite projection of the total angular momentum (i.e. MJ). The theoretical studies also suggested that the most probable thermal relaxation process occurs via the fifth KD and not the highest-energy KD, indicating that the basic molecular design strategy still has room for optimization. It is instructive to compare the SMM properties of the 5 cation to those of the two precursor compounds [(Z5-Ci5Pr5)Dy(BH4)2(THF)] and [(Z5-Cp )Dy(Z5-Ci5Pr5)(BH4)], which have energy barriers of 241(7) cm−1 and 7(1) cm−1, respectively, and closed hysteresis loops at 1.8 K, further highlighting the deleterious effect of the equatorial components of the crystal field provided by the borohydride and THF ligands. The methodology used to synthesize the 5 cation was originally developed as a halide-abstraction route to [(Cpttt)2Dy] [B(C6F5)4] from the reaction of [(Cpttt)2DyCl] with the super-electrophile [(Et3Si)2(m-H)][B(C6F5)4] (Cpttt ¼ 1,2,4-tri-(tert-butyl) cyclopentadienyl).47,48 The same general protocol was used to synthesize the homologous series of dysprosocenium-containing SMMs [(Ci5Pr4R)2Dy][B(C6F5)4], in which R ¼ H, Me, Et or iPr.49 The variation in substituent pattern in the family of dysprosocenium SMMs has yielded important insight into how modifications to the ligand periphery impact on the molecular structure and, in turn, on the SMM properties. Key parameters for all such SMMs are listed in Table 2. The impressive SMM performance of these cations derives from the strong and highly axial crystal field acting on the ground 6 H15/2 multiplet of Dy3+. The blocking temperatures are considerably higher than those found in essentially all other types of SMM based on classical coordination or organometallic ligand environments, and are a consequence of the negligible competing equatorial crystal field, which would otherwise induce mixing between the low-lying crystal field states. The variation in energy barrier and blocking temperature is clearly linked to the Cp-Dy-Cp angle and the Dy–Cp centroid distances, which is, in turn, related to the bulk of the ligands. When the angle approaches 180 and the distances are as short as possible, the key conditions for ‘good’ SMM performance are met. Thus, it is possible for the ligands to be too bulky, as in [(Ci5Pr5)2Dy]+, an arrangement that gives a highly axial crystal field but one that also lengthens the Dy–Cp distances owing to the spatial demands of the substituents. Shorter Dy–Cp distances are possible with smaller substituents, e.g. [(Ci5Pr4H)2Dy]+, but the reduced bulk facilitates bending of the metallocene, which diminishes the axiality. Currently, the optimal combination of bulk is found in the 5 cation, although it is conceivable that further improvement could be made on this qualitative basis by the synthesis of derivatives. That the basic metallocene framework is well-suited to SMM behavior is beyond dispute, but the question as to why such unusual properties occur in dysprosocenium SMMs is only just beginning to be answered. Inevitably, the appealing relationship between energy barrier, blocking temperature and ligand bulk is too simplistic. In terms of an understanding at the microscopic level, state-of-the-art theoretical calculations have shown how interactions between the spin and the surrounding phonon bath play a critical role in the magnetic relaxation.88 Quantitative insight has been obtained through first-principles calculations of first-order spin-phonon couplings as applied to the optical phonons, which can be taken as approximations of molecular vibrations. The first study of this type on a dysprosocenium SMM focused on [(Cpttt)2Dy]+, where it was identified that the CdH oscillators in the ligands vibrate in a manner that initiates the relaxing transition from the first KD to the second KD. Notably, the 5 cation contains no such oscillators. Indeed, the analogous transition in the 5 cation is likely to involve out-of-plane deformations of the C5Me5 ring, a result with the implication that the SMM properties could be improved by targeting the energies of the vibrations by switching to different substituents. A subsequent theoretical study of the vibrational modes of [(Cpttt)2Dy]+ has implied that the unusually good hysteresis properties of dysprosocenium SMMs is related to the relatively inflexible nature of the cyclopentadienyl rings, or their ‘stiffness.’89 In silico design of experimentally unknown complexes, such as the exotic-looking species [(Z5-C5I5)2Dy]+, has focused on the concept of molecular rigidity in order to minimize the extent of spin-phonon coupling and molecular vibrations with energies lower than 1000 cm−1 and, especially, below 500 cm−1.90 If theoretical studies such as these could be translated into experiment, new dysprosium metallocene SMMs with properties that exceed the current state of the art could, in principle, be isolated. The ‘magic’ nature of dysprosium91 in the cationic metallocene SMMs was further highlighted through studies of the dynamic magnetic properties of analogous complexes containing different lanthanides. Terbium is the second most popular lanthanide in

Table 2

Selected structural and magnetic properties for [(CpR)2Dy]+ SMMs.

[(Cp )Dy(C5iPr5)]+ [(Cpttt)2Dy]+ [(C5iPr4H)2Dy]+ [(C5iPr4Me)2Dy]+ [(C5iPr4Et)2Dy]+ [(C5iPr5)2Dy]+ a

Measured in zero applied field. Maximum hysteresis temperature.

b

Cp-Dy-Cp / 

Cp-Dy / ˚A

Ueff / cm–1a

TB / Kb

162.507 152.7 147.2 156.6 161.1 162.1

2.296(1), 2.284(1) 2.316(3) 2.29(1) 2.298(5) 2.302(6) 2.340(7)

1540 1277 1285 1468 1380 1334

80 60 32 72 66 66

228

f-Element Organometallic Single-Molecule Magnets

single-molecule magnetism, having played a pivotal role in developing the field following the seminal discovery of the phthalocyanine (Pc) family of SMMs,92 which subsequently led to many important discoveries in molecular magnetism and molecular spintronics.93,94 In contrast to the high-performance dysprosium versions, a barrier could not be determined for the bent terbocenium cation [(Cpttt)2Tb]+ (Cp-Tb-Cp ¼ 152.2(2) ), and the M(H) hysteresis loops were essentially closed even at 2 K.95 These properties are a consequence of the non-Kramers nature of the Tb3+, which only shows a bistable magnetic ground state when in a coordination environment with strict axial symmetry, as found in, for example, the D4d-symmetric sandwich complex [TbPc2]−.92 A small barrier of Ueff ¼ 39 cm−1 with t0 ¼ 5.1  10−6 s was determined for [(Cpttt)2Ho]+, whereas the erbium, thulium and ytterbium analogues did not show slow relaxation, presumably due to the incompatibility of their prolate 4f electron densities with the dominant axial crystal field.96 In an applied field of 10 kOe, maxima in the w”(n) data were found for [(Cpttt)2Nd]+, and Orbach relaxation with a barrier of Ueff ¼ 51.2 cm−1 and t0 ¼ 9.64  10−8 s were determined. Relaxation in other versions of [(Cpttt)2Ln]+ is affected by weak coordination of the counter-anion via Ln


F interactions, and is thought to be dominated by Raman processes. A general trend among such systems is the involvement of anomalously low Raman exponents (n [red]. However, once the potential sweeps to the reduction potential current starts to flow, [ox] and [red] are equal at Eo0 . As the potential sweeps negative of Eo0 , however, [ox] < [red] at the electrode surface. CV is a method in which the potential is varied over a range like in LSV, but the potential scan is reversed after the Eo0 and toward the initial potential. The intricacies and powerful opportunities afforded by CV will be discussed here and in depth in later sections.36 For those unfamiliar with running CV experiments, we direct the reader to several foundational books and articles which will provide practical guidance for organometallic chemists beginning to utilize electrochemistry. In particular, the books presented by Zoski,40 Bard and Faulkner,36 and Constentin and Savéant29 can provide the reader with a comprehensive guide to the theory, instrumentation, and electrochemical techniques necessary for the investigation of organometallic complexes. Popular educational articles8 and resources41 are available that provide initial information on selecting solvents, electrolytes, electrodes, and use of electrochemical cells to perform a CV experiment. Differential pulse voltammetry (DPV)42 allows interrogation of a system via an approach similar to linear sweep voltammetry, but involving a more complex pulse sequence that minimizes non-Faradaic currents in the final data output. This technique, like CV and LSV, can be used for a variety of mechanistic and kinetic measurement applications.43 DPV can be preferred over CV methods because the series of pulses applied at steadily changing potentials minimizes background from background charging of the electrode.36 Square wave voltammetry43,44 is similar to DPV and utilizes a discontinuous potential change in which its current output is obtained as a symmetrical peak, unlike the current output obtained in linear sweep voltammetry that displays the direct influence of diffusion. Like DPV, square wave is a technique that is useful for detection of very low concentrations of analytes. The use of discontinuous potential in this technique allows measurement of Faradaic currents at the points in which current from double layer charging at the electrode is negligible. Digital computer simulations of voltammetry data33,45,46 can be employed as a complement to experimental cyclic voltammetry to help decipher and quantify coupled electron transfer reactions. Simulations of cyclic voltammograms can be performed by changing variables such as electrode area, solution concentration, scan rate and switching potential, in order to extract kinetic information from voltammograms. Mathematically derived differential equations can be applied to organometallic systems to describe the concentration of a species in solution as a function of time. Other variables that are described by these differential equations include diffusion, convection, and migration; however, understanding these variables can be complex due to heterogenous chemical processes occurring at the electrode surface and homogenous chemical reactions occurring in the bulk solution.47 Some simulation packages available include: ELSIM,48–51 DigiSim,52 DigiElch,53 CVSIM54 and CVPLOT.54 An exhaustive list of packages and description can be found in Digital Simulation in Electrochemistry.47 Low temperature cyclic voltammetry55 is an attractive technique for interrogating mechanistic pathways that create electrogenerated transient intermediate species. Additionally, this technique enables electrochemists to quantify thermodynamic parameters such as enthalpy change and kinetic parameters such as diffusion rate and heterogenous electron transfer rates of electron transfer reactions, and aids in the understanding of intricate coupled chemical reactions. Two cells are suitable for this technique, the dip-type cell that have an elongated cell body suitable for immersion in a coolant and jacketed cell that has a second outer jacket with feed throughs for coolants. Special care needs to be taken to choose an appropriate solvent and supporting electrolyte that can be used at low temperatures; a comprehensive list can be found in Laboratory Techniques in Electroanalytical Chemistry.33 An additional challenge is that solution resistance increases at low temperatures; it is therefore often necessary to electronically compensate for the solution resistance.33 Spectroelectrochemistry56 is a technique that allows for simultaneous electrochemical and spectroscopic interrogation of a compound. Spectroelectrochemistry enables the observation of spectral changes in situ to understand the reactions occurring at an electrode surface. These cells contain an optically transparent electrode (OTE) that are constructed of a conductive material such as platinum, gold, carbon, or a semiconductor material such as doped tin oxide on glass or quartz. Various spectroscopic techniques such as electronic absorption spectroscopy, infrared spectroscopy and Raman spectroscopy can be used in parallel with electrochemical methods.36 The use of chemically modified electrodes,57 in which electrodes are purposefully modified via adsorption, coating or attachment of molecules to the electrode surface, is also a robust area of inquiry relevant to numerous fields. Modified electrodes are studied for their interesting properties in applications such as electrocatalysis, display devices, analytical applications and photoelectrochemical applications. Alternating current (AC) voltammetry is a small-amplitude method that involves the application of sinusoidal wave voltage to an electrochemical cell. This method enables deciphering of contributions to electrochemical behaviors of, separately, the analyte concentration, the identity of solution components, kinetics of charge transfer, and the nature of the double layer capacitance at the electrode/electrolyte interface. This technique can be used primarily for mechanistic studies and allows separation of Faradaic and non-Faradaic current responses.33 Modern electrochemical methods and techniques will be discussed in further detail in the context of organometallic electrochemistry throughout this chapter.

254

Electrochemistry in Organometallic Chemistry

1.09.3

Chemical reactivity at the working electrode surface

1.09.3.1

Electrochemically and chemically reversible electron transfers

Electron transfer can be homogenous or heterogenous in nature. Homogenous electron transfer is the transfer of an electron from one solubilized molecule to another. The movement of an electron from the highest occupied molecular orbital (HOMO) of a donor molecule and the lowest unoccupied molecular orbital (LUMO) of a given acceptor should be thermodynamically favored to take place. The kinetics of homogeneous electron transfer are most commonly discussed in the context of Marcus Theory.58 On the other hand, heterogenous electron transfer typically occurs in molecular electrochemistry, an electron being transferred from a solid electrode to a solubilized molecule in solution that is located near the electrode surface. This type of electron transfer is associated with electrochemical reduction/oxidation in which the driving force for electron transfer to a compound in solution is based on the potential applied to the electrode by the potentiostat. Electron transfer reactions are ubiquitous in organometallic electrochemistry and it is therefore often important to understand how electron transfer to/from an electrode affects the structure and composition of organometallic complexes. Electron transfer reactions are broadly classified as either “reversible” or “irreversible” processes. However, these classifications should be more specific: electron transfer reactions can be classified as either electrochemically (ir)reversible and/or chemically (ir)reversible. Chemical reversibility refers to processes in which the electroactive species are stable and homogenous in their reduced and oxidized forms.8 Electrochemical reversibility speaks to the rate of electron transfer kinetics between the electrode surface and the electroactive species, with fast rates of electron transfer giving rise to electrochemically reversible behavior. Electrochemically reversible systems are referred to as Nernstian, meaning that they obey the Nernst Eq. (1) by virtue of equilibrium concentrations being fully established at the electrode surface.8,36 Processes that are both electrochemically and chemically reversible have a potential separation between cathodic and anodic peaks (DEp) of 57 mV in cyclic voltammetry, a value derived first in the work of Nicholson and Shain for electron transfer rates at the fast limit. The term quasi-reversible refers to electron transfers that appear chemically reversible but have DEp > 57 mV, or for cases where the forward and backward rates are similar but not quite equal (Scheme 1).

Scheme 1 Reversible electron transfer (kf ¼ kb) and quasi-reversible electron transfer (kf  kb).

Reversible electron transfers can be described with the E mechanism notation used by Nicholson and Shain, referring to the involvement of a single electron transfer step.28 The most well-known example of an organometallic complex that engenders an E mechanism upon oxidation and reduction is ferrocene/ferrocenium; ferrocene is ubiquitous in organometallic electrochemistry in part due to its ability to undergo fast electron transfer with minimal redox-induced structural change. Cycling of the oxidation state of the ferrocenium/ferrocene couple results in very little structural distortion, giving rise to fast electron transfer and therefore it is typically electrochemically reversible.59 This has led to the widespread adoption of ferrocenium/ferrocene (Fc+/0) as an internal reference for the reporting of electrochemical potentials for redox events.60 This is because the reversible behavior of the Fc+/0 couple holds true for a range of solvents; there is little to no structural rearrangement upon electron transfer, and therefore there are no significant influences of solvent interactions that would otherwise affect the appearance of the voltammogram. Although ferrocene is traditionally viewed as highly stable and is therefore used as a standard for normalizing potential, readers should be aware that ferrocene can be oxidized to a dication61 or reduced to an anion,62 or degrade from reactions with electrogenerated species. As the reader progresses ahead in this chapter toward the CV and CPE sections, please note that midpoint potentials (E1/2 values) for redox processes may be reported, at various points, versus the ferrcenium/ferrocene couple (Fc+/0) as well as versus SHE, the standard hydrogen electrode. This is because, in order to for a reported potential to be a standard reduction potential (Eo), the conditions of the measurement of the potential must satisfy the Nernst equation (see Eq. 2), wherein the temperature is 25  C, any gases involved are present at 1 atm, and the molarity of the analytes are 1 M.63 These conditions are not always readily achievable or desireable in organometallic chemistry, and thus midpoint potentials measured for various redox processes are often simply reported as E1/2 values vs Fc+/0 or other reference potentials. In order to convert between V vs SHE and V vs Fc+/0, Eo would need to be rigorously determined for ferrocene in water. However, ferrocene is insoluble in water, limiting the ability to rigorously determine the Eo of Fc+/0 or other insoluble redox active species in terms of V vs Fc+/0, limiting the precise and rigorous interconversion between V vs SHE and V vs Fc+/0. Thermodynamic half-cell potentials for processes like proton reduction, oxygen reduction, and water oxidation are typically discussed with referencing of V vs SHE instead of V vs Fc+/0 because such small-molecule reactions often do not have well-defined standard potentials in non-aqueous solvents. We note, however, that the conversion between the two references (SHE and Fc+/0) has been estimated to be Fc+/0 ¼ +400 mV vs SHE.10 Under non-aqueous conditions, the midpoint potential, E1/2, measured using cyclic voltammetry of the Fc+/0 couple can provide an estimate of the formal potential, Eo0 , of the Fc+/0 couple.64 The midpoint potential and formal potential of the ferrocenium/ ferrocene couple are virtually equivalent because of the electrochemically and chemically reversible nature of the redox chemistry of these species. Therefore, a traditional measurement to determine the standard reduction potential of ferrocenium/ferrocene

Electrochemistry in Organometallic Chemistry

255

(carried out so as to satisfy all the requirements of the Nernst equation) in which the potential of a half-cell prepared with a 1:1 mixture of ferrocenium and ferrocene would be compared to a reference potential can be reliably concluded to be comparable to that of the operationally much simpler measurement of the midpoint potential of the Fc+/0 redox system via CV (Scheme 2).

Scheme 2 One-electron reduction of ferrocenium/ferrocene. General E mechanism. Electrode electron transfer reaction.

1.09.3.2

Coupled chemical reactions

Chemical reactions are often coupled to electron transfer, because gain or loss of electrons by a metal complex can cause structural changes and/or changes in the frontier orbitals in the complex. An extensive review of examples of electrochemically induced redox chemistry is beyond the scope of this chapter, but a thorough list of examples is given in a review from Geiger.6 Electronic changes often induce chemical processes due to the instability of formed intermediate species that result from electron transfer. Irreversible electron transfer processes involve a charge transfer that is rate limiting. A non-reversible electron transfer processes involve thermodynamically reversible process as well as a redox-induced chemical reaction (Scheme 3).

Scheme 3 Irreversible electron transfer (kf  kb) and non-reversible electron transfer.

1.09.3.3

EC mechanism

There are multiple mechanisms by which electron transfers and chemical processes can be coupled together at an electrode.28 The simplest of these is the EC mechanism, in which an electron transfer is followed by a first-order or pseudo-first-order reaction that occurs in solution. If the initial electron transfer is fast, the follow-up homogeneous chemical reaction often does not interfere with the electrochemical response in a kinetic sense. In the case of the EC mechanism, the redox-induced chemical reaction is the only rate limiting factor except for diffusion (Scheme 4).29

Scheme 4 General EC mechanism. Electrode electron transfer flowed by a first-order or pseudo-first-order homogenous reaction.

The zone diagrams developed by Savéant are helpful to understand redox reactions that engender a general EC mechanism. Zone diagrams serve to identify the relationship between the rate of diffusion, the kinetic properties of the reduction-induced reaction (as a function of equilibrium constant, K) and the ratio of the coupled chemical reaction rate and diffusion rate (quantified in a dimensionless parameter, l). The scan rate at which the CV is measured is inversely proportional to l. This inverse relationship means that as the scan rate is increased for a CV measurement, the l parameter decreases. When scan rate is high in a reduction, l is small, allowing less time for the reduced form of the organometallic complex to diffuse in solution, and therefore there is less time for the coupled chemical reaction to occur at the electrode surface and influence the electron transfer equilibrium that is measured as current. At faster scan rates, the anodic peak current associated with reoxidation increases until it appears to be quasi-reversible, indicating that the coupled chemical reaction has had insufficient time to occur. Conversely, at low scan rates the coupled chemical reaction appears as an irreversible process, because there is sufficient time for the coupled chemical reaction to occur at the electrode surface, and the back reaction is eliminated.29 Two examples of kinetic zone diagrams are shown in Fig. 1; for a complete set of the diagrams discussed in this chapter please refer to Savéant and Costentin’s Elements of Molecular and Biomolecular Electrochemistry: An Electrochemical Approach To Electron Transfer Chemistry.29 A well-defined organometallic system that demonstrates this behavior with an EC mechanism is the series of [Cp Rh] bis (2-pyridyl)methane complexes studied by our own group.65,66 In this organometallic system, a 1e− quasi-reversible reduction is followed by a second 1e− reduction that is coupled to a rearrangement of the bis(2-pyridyl)methane ligand. Fig. 2 shows the voltammograms of the benzyl (Bn) analog of 1 at 50, 1000, and 2500 mV s−1 (left panel) associated with the reduction-induced

256

Electrochemistry in Organometallic Chemistry

Fig. 1 Examples of zone diagrams. Upper Panel: Kinetic zone diagram of expected shapes of CV response for a catalytic reaction first order in substrate (CA) and catalyst (CP). Lower Panel: Kinetic zone diagram for the ECrECi mechanism. Upper Panel: From Savéant, J.-M. Molecular Catalysis of Electrochemical Reactions. Mechanistic Aspects. Chem. Rev. 2008, 108 (7), 2348–2378 with permission. Lower Panel: From Elias, J. S.; Costentin, C.; Nocera, D. G. Direct Electrochemical P(V) to P(III) Reduction of Phosphine Oxide Facilitated by Triaryl Borates. J. Am. Chem. Soc. 2018, 140, 13711–13718 with permission.

ligand rearrangement shown (right panel). The second reduction at Epc ¼ − 1.40 V vs Fc+/0 exhibits characteristics of an EC mechanism in which there is an expected increase in cathodic current with increasing scan rate coupled with observation of the anodic process at only higher scan rates. As described by Savéant, l is large at slower scan rates allowing ample time for the redox induced ligand rearrangement to occur. At higher scan rates the observation of the reoxidation event in the form of the anodic feature at leading to an electrochemically quasi-reversible process. At higher scan rates the relative increase in anodic current is a result of the small value of l. In other words, the increased scan rate enables the rate of rate of the coupled chemical step to be overtaken, enabling direct observation of redox couple for the species that can otherwise be considered transient. As a complement to this experimental voltammetric study, digital simulations of the experimental data were performed using DigiElch in order to

Electrochemistry in Organometallic Chemistry 257

Fig. 2 Top left panel: cyclic voltammetry of 1 (R ¼ benzyl, Bn) at increasing scan rates. Conditions: electrolyte, 0.1 M TBAPF6 in CH3CN; working electrode, highly oriented pyrolytic graphite. Upper right panel: experimental and simulated voltammograms of 1 (R ¼ methyl, Me). Lower scheme: electrochemical reduction pathway for 1. From Hopkins Leseberg, J. A.; Lionetti, D.; Day, V. W.; Blakemore, J. D. Electrochemical Kinetic Study of [Cp Rh] Complexes Supported by Bis(2-pyridyl)methane Ligands Organometallics. 2021, 40 (2), 266–277 with permission.

258

Electrochemistry in Organometallic Chemistry

measure the first-order rate constant (k+) for the ligand reorientation or “flip” chemical step. Multiple scan-rate dependent voltammograms were simulated to extract, from the behavior of the anodic feature, quantitative values of k+.

1.09.3.4

ECC mechanism

An electrochemical process that involving an electron transfer step followed by two sequential chemical reactions is referred to as an ECC mechanism (Scheme 5). One of the most common ECC mechanistic scenarios results from the generation of an organometallic radical, followed by the loss of a ligand, and then dimerization to give the final product. This can occur by 1e− oxidation or reduction of a substrate, which ultimately results in a ligand or metal centered radical. Radicals are notoriously reactive species, and once generated electrochemically, often go on to react with another species present. A notable organometallic complex that undergoes 1e− reduction to form a metal centered radical is Mn(CO)3Br(Hbpy), a well-known CO2 reduction catalyst. The electrochemical and electrocatalytic behavior of Mn(CO)3Br(Hbpy) was first established by Deronzier, Chardon-Noblat, and co-workers using CV (see Fig. 3).67

Scheme 5 General ECC mechanism. Electrode electron transfer followed by two first-order or pseudo-first-order homogenous reactions.

Fig. 3 CV of Mn(CO)3Br(Hbpy) in 0.1 M TBAPF6/MeCN electrolyte at 100 mV s−1 (left). Accompanying chemical scheme for the sequential irreversible reductions and more positive oxidation associated with Mn(CO)3Br(Hbpy) (right). Modified from Bourrez, M.; Molton, F.; Chardon-Noblat, S. Deronzier,[Mn(bipyridyl)(CO)3Br]: An Abundant Metal Carbonyl Complex as Efficient Electrocatalyst for CO2 Reduction Angew. Chem. Int. Ed. 2011, 50 (42), 9903–9906.

At first glance, the cyclic voltammogram of Mn(CO)3Br(Hbpy) seems to be quite complicated; this is attributable to the multiple chemical reactions involved in the redox behavior of this complex. Scanning toward more negative potentials, two sequential and essentially irreversible reductions are observed with cathodic peak potentials of −1.61 V vs Fc+/0 and −1.83 V vs Fc+/0, followed by an oxidation at more positive potentials with an anodic peak potential at −0.61 V vs Fc+/0. Based on mechanistic work with this complex, the first irreversible reduction of Mn(CO)3Br(Hbpy) can be reliably assigned to result in a 19e− complex.68 This 19e− complex then loses Br− to generate a 17e− species with a Mn-centered radical. Two of these Mn-centered radicals can recombine to dimerize and produce Mn2(CO)6(Hbpy)2 (Mn-Mn dimer) in an overall ECC-type mechanism. The second reduction feature in the CV is attributed to the reduction of the Mn-Mn dimer, which cleaves the metal-metal bond to generate the 18e− complex, [Mn(CO)3(Hbpy)]− in an overall EC process. Lastly, scanning to more positive potentials, [Mn(CO)3(Hbpy)]− is re-oxidized to form the starting complex, Mn(CO)3Br(Hbpy).

1.09.3.5

CE mechanism

The CE mechanism involves a rapid, reversible first-order or pseudo-first-order chemical reaction followed by an electron transfer. For CE mechanisms the electroactive starting material (A) is the product of chemical conversion of the starting material (species C in Scheme 6) in the chemical reaction labeled as step C.

Electrochemistry in Organometallic Chemistry

259

Scheme 6 General CE mechanism. A first/pseudo first-order homogenous reaction preceding an electrode electron transfer.

CE mechanisms in which the pre-equilibrium is sufficiently fast to not influence the electrochemical response kinetically can be described using the equilibrium constant (K) and kinetic parameter l.29 In the case of one-electron reversible waves that engender a CE mechanism when the equilibrium constant is large, the couple appears reversible regardless of l. However, when K is small, the appearance of the CV becomes more complex and depends on the value of l. The preceding reaction (C step) influences the electrochemical response (E step) when kf is small. The CV response is controlled by the rate at which the species C is converted to the electroactive species A (kf). The kinetics and thermodynamics of the conversion of C to A can influence the height and shape of the CV. As previously stated, the inverse relationship between scan rate and l means that as the scan rate is increased for a CV measurement, the l parameter decreases. The case of CE with a small l engenders (fast scan rate and/or slow reaction) a reversible wave in which the height is influenced by the equilibrium constant. If the equilibrium constant is proportional to the peak heights of the CV, meaning for a CE case with a small K and small l value there will be a reversible wave with small peak current. As l increases as a result of slower scan rate the anodic wave takes the form of a plateauing wave which has a plateauing current independent of scan rate. An even slower scan rate and larger l parameter results in a CV that gradually increases in reversibility until reaching a quasi-reversible wave.29 A distinct CE mechanism was observed in the titration of the solvent rhodium complex [Cp Rh(PQN)NCCH3]2+ ((PQN) ¼ (diphenylphosphino)quinoline) with tetrabutylammonium chloride (TBACl).69 This titration was performed to understand the role of halide ligands in influencing the redox properties of [Cp Rh] complexes supported by bidentate chelates. The CV of [Cp Rh(PQN)NCCH3]2+ reveals two distinct quasi-reversible one-electron couples (RhIII/II and RhII/I, −0.93 V and −1.16 V vs Fc+/ 0 respectively). Increasing the concentration of TBACl to [Cp Rh(PQN)NCCH3]2+ clearly results in the diminution of the two original waves and gives rise to a new single 2e− wave (RhIII/II ¼ −1.19 V vs Fc+/0) that corresponds to the electrochemical profile of [Cp Rh(PQN)Cl]+; this was verified when this species was isolated through chemical synthesis and studied independently. This behavior confirms the rapid chemical conversion (C step) of [Cp Rh(PQN)NCCH3]2+ to [Cp Rh(PQN)Cl]+ by displacement of the ligated CH3CN with chloride followed by an electron transfer to reduce the rhodium(III) complex (E step) giving rise to an EC mechanism observed in Fig. 4.70

Fig. 4 Electrochemical response of [Cp Rh(PQN)NCCH3]2+ in CH3CN upon increasing additions of tetrabutylammonium chloride. Growth of the reduction process with E1/2 ¼ −1.19 V indicates coordination of chloride to [Cp Rh(PQN)NCCH3]2+ by displacement of bound CH3CN, generating [Cp Rh(PQN)Cl]+. From Hopkins, J. A.; Lionetti, D.; Day, V. W.; Blakemore, J. D. Chemical and Electrochemical Properties of [Cp Rh] Complexes Supported by a Hybrid Phosphine-Imine Ligand Organometallics. Organometallics 2019, 38, 1300–1310 with permission.

260

1.09.3.6

Electrochemistry in Organometallic Chemistry

ECE mechanism

The ECE mechanism is the case in which an electroactive species undergoes an electron transfer reaction followed by a coupled chemical reaction, like that of an EC reaction. However, the intermediate formed by the chemical step then undergoes a second electron transfer (Scheme 7). An ECE reduction occurs when the intermediate C formed has a standard reduction potential that is more positive than that of the standard reduction potential for the conversion of the initial electroactive species A to generate B. Each molecule of B is chemically converted to C, and following the chemical conversion, C is instantaneously reduced by the electrode to generate the final species, D. This instantaneous reduction of C means that the ECE mechanism results in the appearance of a “2e− wave” which is often misinterpreted as arising from a truly simultaneous transfer of 2e−. However, virtually all ECE reductions actually involve two separate 1e− reduction steps. Like many of the other mechanisms discussed here, the appearance of the CV can sometimes be influenced by varying the scan rate. In the ECE case, the l kinetic parameter is interrelated with the rate of the chemical reaction that controls conversion of B to C. In the ECE case, at faster scan rates where l is small, a quasi-reversible wave is predicted to be observed, corresponding to the one-electron process for reduction of A to B and re-oxidation of B to A. As the scan rate is decreased, however, l increases, giving rise to irreversibility with respect to the original A/B wave and increase of the cathodic peak current, attributable to time being allowed for conversion of B to C and subsequent reduction of C to D. At slow scan rates the CV should then have the expected twofold increase in the magnitude of the cathodic peak current. In certain systems which have appropriate kinetics of the chemical step associated with conversion of B to C, accumulation of the intermediate C at the electrode during scans can give rise to trace crossing; such behavior is attributable to the greater quantity of the intermediate C present at the surface during the reverse scan than the cathodic scan, giving rise to cathodic current during the anodic scan.29 However, at typical electrodes and scan rates of relevance to molecular organometallic chemistry, this trace crossing associated with the ECE mechanism is not commonly observed and thus the absence of the behavior should not be taken as evidence against net 2e−, ECE-type behavior.

Scheme 7 General ECE mechanism (reduction shown). Electrode electron transfer flowed by a first/pseudo first-order homogenous reaction which undergoes a second electron transfer.

A well-defined family of Cp Rh complexes illustrates the ECE mechanism engendering two sequential 1e− reductions.66,71–74 Electrochemical studies on the [Cp Rh(PQN)L]n+ (PQN ¼ 8-(diphenylphosphino)quinoline) complex exhibits interesting activity with substitution of the monodentate ligand (Fig. 5). The [Cp RhIII(PQN)Cl]+ analog undergoes an initial 1e− reduction generates a transient [Cp RhII(PQN)Cl]0 following the initial E step, the Rh(II) complex undergoes a chemical loss of the monodentate Cl− to form the intermediate [Cp RhII(PQN)]+, which is immediately followed by a second 1e− reduction to generate [Cp RhI(PQN)]0. This ECE mechanism gives rise to a CV with the appearance of a quasi-reversible 2e− reduction. In the absence of chloride, however, the reduction of the solvento species, [Cp RhIII(PQN)NCCH3]2+ takes place at a potential more positive than that for

Fig. 5 CV of [Cp RhIII(PQN)Cl]+ (upper panel) [Cp RhIII(PQN)NCCH3]2+ (middle panel) and [Cp RhI(PQN)]0 (lower panel). From Hopkins, J. A.; Lionetti, D.; Day, V. W.; Blakemore, J. D. Chemical and Electrochemical Properties of [Cp Rh] Complexes Supported by a Hybrid Phosphine-Imine Ligand Organometallics. Organometallics 2019, 38, 1300–1310 with permission.

Electrochemistry in Organometallic Chemistry

261

[Cp RhIII(PQN)Cl]+ due to its dicationic nature. The shift in this potential is a result of the chemical reaction step (loss of CH3CN rather than Cl−) being insufficient to cause the transient [Cp RhII(PQN)NCCH3]+ complex to be reduced at a more positive potential compared to the [Cp RhIII(PQN)NCCH3]2+ complex. Due to this phenomenon two discrete reductions are observable and results in the metastable [Cp RhII(PQN)NCCH3]+ complex generated near the electrode. The oxidation of the [Cp RhI(PQN)]0 analog results in a very similar CV profile due to the generation of the [Cp RhIII(PQN)NCCH3]2+ in CH3CN solvent. The third reduction in each CV is attributed to the reduction of the PQN ligand.74

1.09.3.7

Substitution mechanism

In inorganic and organometallic chemistry intermolecular substitution reactions typically occur by associative or dissociative mechanisms. In an associative mechanism, a ligand forms a bond to the metal center prior to the ejection of the exiting ligand. Conversely, in a dissociative mechanism, a ligand is initially expelled from the metal center to generate space for an incoming ligand to bind. Electrochemical oxidation or reduction has been shown to accelerate some intermolecular organometallic substitution reactions, and therefore gives useful mechanistic insights. In an example from the work of Kochi and co-workers, organometallic intermolecular ligand substitution is enhanced and observed using electrochemical methods.75 This series of reactions, monitored by CV, demonstrate that the acetonitrile ligand bound to (Z5-MeCp)Mn(CO)2(NCMe) (Mn-NCMe) may be substituted with various alkyl and arylphosphines when an oxidative bias is applied (see Fig. 6). The initial CV of Mn-NCMe displays a well-behaved quasi-reversible 1e− redox couple at 0.22 V vs Fc+/0. However, when 1 equiv. of triphenylphosphine (PPh3) is added to the system, the quasi-reversibility of the initial redox couple ceases, but the 1e− oxidation of Mn-NCMe to the 17e− species [Mn-NCMe]+ persists and a new redox couple begins to grow in at 0.55 V vs Fc+/0. This behavior suggests that once [Mn-NCMe] is oxidized to [Mn-NCMe]+, substitution of the NCMe for PPh3 is facile

Fig. 6 (A) Initial CV of (Z5-MeCp)Mn(CO)2(NCMe). (B and C) Potentiometric titration of (Z5-MeCp)Mn(CO)2(NCMe) (R) with PPh3 to generate (Z5-MeCp)Mn(CO)2(PPh3) (P) under electrochemical conditions. (D) CV of isolated (Z5-MeCp)Mn(CO)2(PPh3). CVs are taken at 200 mV s−1 with MeCN solvent at 0.1 M TEAP electrolyte. From Hershberger, J. W.; Klingler, R. J.; Kochi, J. K., Kinetics, thermodynamics, and mechanism of the radical-chain process for ligand substitution of metal carbonyls. J. Am. Chem. Soc. 1983, 105, 61 with permission.

262

Electrochemistry in Organometallic Chemistry

and results in the rapid consumption of [Mn-NCMe]+ to generate the intermolecular substitution product (Z5-MeCp)Mn(CO)2 (PPh3) (Mn-PPh3). As increasing equivalents of PPh3 are added to the electrochemical cell, [Mn-NCMe]+ is consumed more rapidly until only the current corresponding to the substitution product is observed during CV. To confirm that it is indeed (MeCp) Mn(CO)2(PPh3) being generated, an authentic sample of (MeCp)Mn(CO)2(PPh3) was found to be oxidized at the latter potential. This exemplifies a class of substitution reactions that take place when inner-sphere ligands become labile upon oxidation or reduction of different organometallic complexes.

1.09.3.8

DISP mechanism

The DISP mechanism, or disproportionation mechanism, is similar to that of the ECE mechanism previously discussed, but involves a second disproportionation pathway following the non-electrode-based generation of the intermediate C (Scheme 8). For DISP mechanisms the second electron is transferred from B to C rather than from the electrode to C as in an ECE mechanism. In the ECE mechanism the EoC/D is at a more positive potential than EoA/B, implying that the disproportionation reaction KD is favorable. The two electron transfers engendered in the ECE mechanism are both fast, therefore it is assumed that the thermodynamically favorable disproportionation reaction is fast. The chemical conversion of B to C is the rate determining step and l behaves in the same manner as in ECE mechanism. As discussed for the ECE mechanism at faster scan rates l is low, which gives rise to a quasi-reversible wave. As the scan rate is gradually decreased, l increases giving rise to increasing irreversibility and increase of the cathodic peak current and slow scan rates the CV reaches complete irreversibility and has appearance of a 2e− reduction in magnitude of the peak current. Under pure kinetic conditions the ECE and DISP mechanisms are indistinguishable by CV, but when pure kinetic conditions are not achieved it is possible to distinguish between ECE and DISP mechanisms. Non-pure kinetic conditions in DISP do not lead to crossing as seen in the ECE case discussed above. The absence of trace crossing in the DISP mechanism is due to the slow reaction of B to C, consequently the intermediate C is formed far from the electrode surface which allows ample time for the disproportionation to occur before C reaches the electrode. This prevents the accumulation of C at the electrode surface during scans, and therefore there is no possible contribution of cathodic current during the anodic scan as is possible during the ECE-type mechanism.29

Scheme 8 DISP mechanism. Electrode electron transfer followed by a first-order or pseudo-first-order homogenous reaction in which a second electron is transferred from species B to C.

An example of an organometallic system with an ECE-DISP mechanism is the series of substituted rhodium(III) porphyrin complexes studied by Savéant et al. (Fig. 7).76 This rhodium(III) porphyrin family engenders chemical irreversibility that is

Fig. 7 Reduction scheme of Rh(III) porphryin complexes in which disproportionation occurs with the reduction to Rh(II)L2. Disproportionation scheme of Rh(II) L2 shown below. (Left) Cyclic voltammetry of rhodium(III) tetraphenylporphyrin (TPP) and rhodium(III) octaethylporphyrin (OEP) (Right). From Grass, V.; Lexa, D.; Momenteau, M.; Savéant, J.-M. Reductive Electrochemistry of Rhodium Porphyrins. Disproportionation of Intermediary Oxidation States. J. Am. Chem. Soc. 1997, 119 (15), 3536–3542 with permission.

Electrochemistry in Organometallic Chemistry

263

attributed to the disproportionation of the rhodium(II) complexes caused by ligand exchange reactions. During reduction + Rh(III) L2 is reduced by 1e− to the corresponding Rh(II)L2 followed by the loss of one L ligand. The loss of L generates a second Rh(II)L which is reduced at a more positive potential than the initial Rh(III) complex by 1e−. This electrochemical reduction thus involves a disproportionation step because the second electron step occurs concurrently with electrode reduction.

1.09.4

Using electrochemistry to explore chemical reactivity with stoichiometric redox reagents

1.09.4.1

Synthesis via controlled potential electrolysis

Controlled potential electrolysis (CPE) is a well-established electrochemical method used for synthesis of electron transfer products.10,77–82 There are advantages and disadvantages to this electrochemical synthetic method. CPE allows a wide range of precise potentials accessible for reduction/oxidation of organometallic compounds with minimal side reactivity that leads to byproducts. However, CPE can be problematic because of the need for a large amount of electrolyte salt in solution, which can be difficult to separate from the product.10

1.09.4.2

Synthesis via chemical redox reagents

An alternative method for synthesis of electron transfer products is use of chemical redox reagents, chosen based on CV to determine the reduction potential of the electroactive species (Fig. 8). The utilization of a homogeneous reductant/oxidant eliminates the necessity for a large excess of supporting electrolyte salt that must be separated. Another advantage of chemical methods is the fast time scale and large reaction scale compared to electrochemical methods. Chemical reduction and oxidation can be performed at low temperatures to reduce side reactivity and improve selectivity. Although it should be noted that low temperature electrolysis experiments are possible, it is necessary to be strategic when selecting the solvent and supporting electrolyte for work at low temperature. A list of appropriate solvents and electrolyte systems for work at reduced temperatures can be found in Kissinger and Heineman’s Laboratory Techniques in Electroanalytical Chemistry.83 Another notable difference between chemical and electrochemical synthetic methods is the former can be done in non-polar solvents. This may be advantageous for precipitation of products, or for use of a non-coordinating solvent to prevent ligand exchange/displacement.10 However, the use of chemical redox reagents have notable disadvantages.10 Chemical reagents have a fixed reduction/oxidation potential, which can be disadvantageous for target complexes for which there is no good match of a chemical redox agent with chemical compatibility and the correct potential, whereas electrochemical methods enable precise choice of potential. Chemical reagents also introduce a redox byproduct, which may give undesired reactivity. Particularly problematic are reagents that give inner-sphere electron transfer leading to installation of ligands such as NO when using the oxidant [NO]+ or chloride derived from trityl chloride.10,84 In choosing a redox reagent for chemical reduction or oxidation, one must take solubility into account as well as considering the reduction potential of the electroactive species and fixed potential of the redox reagent. It is important to note that for a Nernstian system (see Section 1.09.3.1) a 99% reaction completion for a 1e− transfer requires a reducing agent to have Eo 118 mV negative of the Eo and an oxidizing agent to have Eo 118 mV positive of the Eo (Fig. 8).10

1.09.4.3

Literature examples of combined use of cyclic voltammetry and chemical redox reagents

An example of utilizing chemical reducing agents as a supporting complement to cyclic voltammetry is a titanium polymerization catalyst with added AlEt3.85 This work features cyclic voltametric studies of the titanium catalyst Ti-1 shown in Fig. 9 in which the complex undergoes a single irreversible 1e− reduction at −2.12 V vs Fc+/0. The cyclic voltammogram of Ti-1 in the presence of AlEt3 changes the appearance to include an oxidative feature at Epa at −1.32 V vs Fc+/0. Chemical reduction of Ti-1 in the presence of AlEt3 yields bimetallic complex Ti-2 (Fig. 9). The anodic feature at Epa at −1.32 V vs Fc+/0 in the electrochemistry corresponds to the oxidation of Ti-2. This assignment is supported by matching spectral profiles of the spectroelectrochemical data of the parallel addition of AlEt3 to 1 and the UV-vis data of the chemically prepared complex Ti-1. Cyclic voltammetry published by Kubiak et al. of a Mn(bpy-tBu)(CO)3Br complex exhibits two sequential, irreversible 1e− reductions at −1.39 V and −1.57 V vs SCE. The first reduction is assigned to metal centered with concomitant loss of Br− resulting in the formation of a Mn-Mn dimer68 and the second reduction is attributed to formation of [Mn(bpy-tBu)(CO)3]−.86 The feature at −0.30 V vs SCE is assigned to the oxidative cleavage of the Mn–Mn bond. The identity of the second reduction at −1.57 V vs SCE was interrogated by reduction of Mn(bpy-tBu)(CO)3Br by two equivalents of potassium graphite (KC8, Eo  −3.1 V vs Fc+/0 in THF)10 in the presence of 18-crown-6 to yield the doubly reduced [Mn(bpy-tBu)(CO)3][K(18-crown-6)(THF)] (Fig. 10). Note the use of a much stronger reducing agent than formally needed for the reduction, which is common. The solid-state structure of the doubly reduced form was used to the assignment of the second 1e− reduction as the formation of [Mn(bpy-tBu)(CO)3]−. Another literature example of combining cyclic voltammetry methods and chemical reductants comes from a study of dimeric tungsten species.87 The CV of complex W-1 in Fig. 11 exhibits an irreversible oxidation at −2.21 V vs Fc+/0 followed by a corresponding reduction feature indicating an EC process in which a new tungsten species is formed in situ. To interrogate the identity of this species formed the starting material W-1 was chemically reduced by potassium graphite (KC8, E o −3.1 V vs Fc+/0 in THF)10 and it was possible to isolate and crystallographically characterize W-3. The CV of isolated W-3 reveals a quasi-reversible couple at −2.29 V vs Fc+/0 that matches the oxidation feature in the CV of W-1. The paper hypothesizes that the in situ formation of W-3 is most likely a result of disproportionation reactions.

264

Electrochemistry in Organometallic Chemistry

Fig. 8 List of chemical redox agents as summarized by Connelly and Geiger in 1996. Chemical oxidants (left) chemical reductants (right). From Connelly, N. G.; Geiger, W. E. Chemical Redox Agents for Organometallic Chemistry. Chem. Rev. 1996, 96 (2), 877–910 with permission.

A family of [Cp Rh(Me2dpma)] (Me2dpma ¼ dimethylbis(2-pyridyl)methane) complexes were studied for their redox properties as they relate to the [Cp Rh(bpy)] system (Fig. 12).67 The cyclic voltammetry of 1-NCCH3 displays an initial quasi-reversible 1e− reduction from Rh(III) to Rh(II) at −0.85 V vs Fc+/0, and continuing to scan cathodically a second irreversible 1e− reduction occurs at −1.50 vs Fc+/0. The irreversibility of this reduction suggests an EC process. To probe the identity of the species formed upon

Electrochemistry in Organometallic Chemistry

265

Fig. 9 Cyclic voltammetry data for Ti-1 (red trace upper panel) and cyclic voltammetry of Ti-1 with the addition of 6 equivalents of AlEt3 (blue lower panel) (left panel.) Structures of compound Ti-1 and Ti-2 (top panel) and solid-state structures of complex Ti-1 and Ti-2. From Barr, J. L.; Kumar, A.; Lionetti, D.; Cruz, C. A.; Blakemore, J. D. Understanding the Roles of Triethylaluminum in Phosphinimide-Supported Titanium Catalyst Systems for Ethylene Polymerization. Organometallics 2019, 38 (9), 2150–2155 with permission.

Fig. 10 Cyclic voltammogram of Mn(bpy-tBu)(CO)3Br (left) Solid state structure of chemically prepared doubly reduced [Mn(bpy-tBu)(CO)3][K(18-crown6) (THF)] (right). From Smieja, J. M.; Sampson, M. D.; Grice, K. A.; Benson, E. E.; Froehlich, J. D.; Kubiak, C. P. Manganese as a Substitute for Rhenium in CO2 Reduction Catalysts: The Importance of Acids. Inorg. Chem. 2013, 52 (5), 2484–2491 with permission.

electrochemical reduction the 1-Cl and 1-NCCH3 complexes were treated with chemical reductants. Treatment of 1-Cl with cobaltocene (CoCp2, E1/2 ¼ −1.30 V vs Fc+/0)10 yielded the expected Rh(II) reduction product with minimal chemical change as indicated by the 1e− quasi-reversible couple. The 1-NCCH3 complex was used to interrogate the EC process of the doubly reduced product to avoid side reactivity seen with the 2e− chemical reduction of 1-Cl. Treatment of 1-NCCH3 with a stronger reducing agent, sodium amalgam (Na(Hg), Eo ¼ − 2.4 V vs Fc+/0),10 generated the doubly reduced Rh(I) product 3. Upon the addition of the second electron there is a significant ligand rearrangement in which one pyridine moiety “flips” on the bidentate Me2dpma ligand. This assignment of the doubly reduced species identifies the chemical step upon the second reduction. The origin of this rearrangement is attributed to the ability of the Me2dpma ligand to stabilize the low-valent Rh(I) center with strong p-backbonding by way of facial coordination of the pyridine moiety.

266

Electrochemistry in Organometallic Chemistry

Fig. 11 Cyclic voltammogram of W-1 and W-3 (left) chemical reduction of W-1 with KC8 to W-3 (upper right) solid state structure of W-3 (lower right). From Chapovetsky, A.; Patel, P.; Liu, C.; Sattelberger, A. P.; Kaphan, D. M.; Delferro, M. Electrochemical Investigation of Low-Valent Multiply M M Bonded Group VI Dimers: A Standard Chemical Reduction Leads to an Unexpected Product. Organometallics 2020, 39 (24), 4430–4436 with permission.

Fig. 12 Cyclic voltammogram of 1-NCCH3 (left) reduction scheme of 1-Cl with CoCp2 to 2 and 1-NCCH3 with Na(Hg) to 3 (right). Modified from Lionetti, D.; Day, V. W.; Lassalle-Kaiser, B.; Blakemore, J. D. Chem. Commun. 2018, 54 (14), 1694–1697.

1.09.5

Fundamental concepts of organometallic electrocatalysis

1.09.5.1

Motivation for studying organometallic electrocatalysts in the context of energy science

Energy consumption is anticipated to continue increasing as the Earth’s population grows, motivating work to develop new renewable energy sources and technologies that can produce fuels and chemicals with reneable energy. The most abundant renewable energy source available is sunlight, which provides enough energy (1.2  1014 kWh) in 80 min to sustain the global power demand for over a year.88–90 While there is plenty of potential energy available in the form of sunlight, utilizing this energy is challenging due to its intermittency, variable intensity, and uneven distribution across the surface of the Earth. Development of model molecular electrocatalysts capable of converting electrical energy into stored fuels and chemicals can provide fundamental insights into how such catalysts convert energy. Considering that harnessing renewable energy sources is an important goal of contemporary chemistry, organometallic chemistry has been heavily utilized in recent years for development of new electrocatalysts. Electrocatalysts are typically classified as heterogeneous or homogeneous, and there are many relevant examples of studying both classes of catalysts using electrochemical techniques, particularly cyclic voltammetry (CV) and controlled potential electrolysis (CPE). In this section, priority will be given to discussion of homogeneous molecular catalysts because these are (i) more relevant to organometallic chemistry and (ii) mechanistic discussions are more certain for cases in which catalyst (or precatalyst) structures are well defined. CV and CPE experiments are regularly used to interrogate candidate organometallic systems for oxidative or reductive electrocatalytic applications. This section will provide a brief overview of catalytic CV and CPE experiments to investigate organometallic complexes as potential electrocatalysts. CV techniques can also be employed to assist in distinguishing heterogeneous versus homogenous catalysis. CV and controlled potential electrolysis experiments are often used synergistically to extract important thermodynamic, kinetic, and mechanistic information about these electrocatalytic transformations. When these electrocatalytic experiments are combined with product detection, Faradaic efficiency (FE), turnover number (TON), and turnover frequency (TOF) parameters can be extracted. In this section, we will discuss some of the most popular areas of organometallic electrocatalysis,

Electrochemistry in Organometallic Chemistry

267

including literature examples that utilize organometallic complexes capable of electrocatalytic water (WO) and ammonia oxidation (AO), and proton (HER), carbon dioxide (CO2R), dinitrogen reduction reactions (N2RR), and electroorganic transformations (EOT).

1.09.5.2 1.09.5.2.1

Investigating organometallic electrocatalysis using CV Electrocatalytic parameters and thermodynamic overpotential determined by CV

For catalytic CV experiments, the observations and parameters associated with the generation of a catalytic response are well detailed in the viewpoint provided by Appel and Helm.13 Their work provides insight and information about the parameters gleaned from a set of catalytic CV experiments regarding electrocatalytic overpotential. Specifically, parameters associated with the electrocatalytic response such as peak catalytic current (icat), half the peak catalytic current (icat/2), and catalytic potential (Ecat/2) are used in this area to determine the thermodynamic overpotential of a molecular catalyst (see Fig. 13). The icat is the maximum current that flows in a catalytic CV experiment; for a plateauing catalytic wave, this is assigned as the maximum current anywhere in the plateauing region of the wave, while for a non-plateauing wave, this is assigned as the absolute maximum current of the wave. icat/2 is defined as half the catalytic current, and the potential at which this current is observed is termed Ecat/2. Selection of Ecat/2 in this manner ensures that there is minimal variation in the catalytic rate when determining a catalytic potential. The information provided by these CV experiments also provides essential information about the overpotential for a catalytic reaction, which is the driving force beyond the thermodynamic minimum needed to carry out a particular chemical transformation. The thermodynamic overpotential is most accurate when the Nernst equation is satisfied and can be determined using the following equations: EX ¼ EoX +

RT ½HB + Š ln nF ½BŠ

EX ¼ EoX − 0:05916 V  pH

(4) (5)

Notably, overpotential can be determined in aqueous or non-aqueous solvents, but the thermodynamic potential of a catalytic reaction at standard conditions (defined as the equilibrium potential (EoX)), pH of the system, and the Ecat/2 must be defined. Acetonitrile is generally the preferred solvent because EoX can be determined91 and there is well defined pKa scale for this purpose.92–94 As an example, if the reader is interested in proton reduction to dihydrogen using an acid such as protonated dimethylformamide, [DMFH]+, as the substrate, a series of cyclic voltammograms would be run at standard-state conditions, with a buffered solution of DMF/[DMFH]+, and under 1 atm of H2. Under these conditions, icat, icat/2, and Ecat/2 can be determined to calculate the thermodynamic overpotential. Though the observation of a catalytic response during a CV experiment does not depend on satisfying the Nernst equation for a particular transformation, an accurate calculation of the thermodynamic overpotential can only extracted when this is possible. A relevant example uses the organometallic complex [Cp RhCl(tBubpy)]+ (where Cp is Z5-(1,2,3,4,5-pentamethylcyclopentadienyl) and tBubpy is 4,40 -bis-tertbutyl-2,20 -bipyridine) as a proton reduction catalyst for the production of dihydrogen (see Fig. 14).73 In the

Fig. 13 Simulated catalytic wave for a molecular catalyst showing the selection of icat, icat/2, and Ecat/2 (left). Experimental CV illustrating determination of Ecat/2 (and overpotential) for H2 production (right). From Appel, A. M.; Helm, M. L. Determining the Overpotential for a Molecular Electrocatalyst. ACS Catal. 2014, 4 (2), 630–633 with permission.

268

Electrochemistry in Organometallic Chemistry

Fig. 14 Pathway for HER using [Cp Rh(Rbpy)] showing the energy loading and energy storage steps (left). Cyclic voltammograms with [Cp RhCl(tBubpy)]+, 1–5 equiv. of PhNH2/[PhNH3]OTf, and 1 atm of H2 at a scan rate of 100 mV s−1 (right). Modified from Henke, W. C.; Lionetti, D.; Moore, W. N. G.; Hopkins, J. A.; Day, V. W.; Blakemore, J. D. Ligand Substituents Govern the Efficiency and Mechanistic Path of Hydrogen Production with [Cp Rh] Catalysts. ChemSusChem 2017, 10 (22), 4589–4598.

absence of acid, the CV of [Cp RhCl(tBubpy)]+ exhibits a quasi-reversible redox event centered at −1.25 V vs Fc+/0 in MeCN. When acid is added, the quasi-reversible behavior of the RhIII/RhI redox couple ceases, and a catalytic response is observed. The loss of reversibility here indicates that the acid is reacting with the reduced RhI species and the resulting current enhancement is consistent with catalytic behavior. The current enhancement results from the RhIII starting material being reduced to RhI, which then interacts with the acid substrate, evolving dihydrogen, and subsequently regenerating the RhIII starting complex. Since the potential at the electrode surface is still sufficiently reducing, this process continues as long as the potential of the electrode remains at values sufficient to drive the reduction processes involved in the chemistry. With these CVs in hand, the icat (−4.16 mA cm−2), icat/2 (−2.08 16 mA cm−2), Ecat/2 (−1.33 V), and overpotential (0.569 V) were determined for the catalytic process. Taking these parameters into account, the CPE experiments were carried out under identical conditions with polarization at −1.36 V vs Fc+/0 over the course of 90 min. Other [Cp RhCl(Rbpy)]+ (R ¼ 4,4-bis-substituted-2,20 -bipyridyl, H and CF3) complexes have also shown this catalytic behavior in the presence of a proton source. These catalytic CV experiments are not limited to proton reduction and have been carried out in the exploration of catalytic applications involving WO, AO, HER, CO2R, and N2RR.

1.09.5.2.2

Kinetic considerations in electrocatalysis: Foot of the wave analysis

To extract kinetic and latent mechanistic information from CV experiments, one employs a method pioneered by Constentin, Savéant, and co-workers known as foot-of-the-wave analysis (FOWA).95 This method simplifies the analysis because the earliest part of the catalytic wave inevitably has pseudo-first order kinetics because the substrate will be in vast excess compared to the catalyst.96 FOWA minimizes the effects of secondary phenomena such as catalyst inhibition, decomposition, or saturation, and maximizes the information provided by the current that flows. Under ideal conditions, an observed rate constant can be computed using:  12 2:24 RT icat Fv 2kobs CA  ¼ F ip 1 + exp RT ðE − Eo Þ

(6)

where icat is the peak catalytic current, ip is the peak current of the redox process of the catalyst in the absence of substrate, R is the ideal gas constant, F is Faraday’s constant, T is temperature, n is the scan rate, kobs is the observed rate constant, CA is the initial concentration of substrate, E is the present potential and Eo ¼ Ecat/2 which is the thermodynamic potential for the catalytic wave. Plotting icat/ip as a function of 1/{1 + exp.[(F/RT)(E −Eo)]} produces a line with a slope equal to 2.24((RT/Fn)2kobsCA)1/2, that can be used to extract the observed rate of catalysis (see Fig. 15). These experiments can be repeated at various substrate concentrations, and at several different scan rates, to extract an average rate constant for the desired electrocatalytic reaction. The average kobs is computed from a linear regression of the scan rate dependent data with a fixed slope of zero, since the intrinsic chemical kinetics should not depend on scan rate. This kinetic technique has become a recognized approach to determining substrate order and observed rates when mapping out the mechanistic details of electrocatalysts. FOWA has been used in various electrocatalytic kinetic studies for WO,97 AO,98 HER,99 CO2R,100 and N2RR.101

Electrochemistry in Organometallic Chemistry

269

Fig. 15 An example FOWA plot of icat/ip as a function of 1/{1 + exp.[(F/RT)(E−Eo)]}. The pseudo-first order rate constant is extracted from the initial linear region of the plot. From Clark, M. L.; Cheung, P. L.; Lessio, M.; Carter, E. A.; Kubiak, C. P. Kinetic and Mechanistic Effects of Bipyridine (bpy) Substituent, Labile Ligand, and Brønsted Acid on Electrocatalytic CO2 Reduction by Re(bpy) Complexes. ACS Catal. 2018, 8 (3), 2021–2029 with permission.

1.09.5.2.3

Distinguishing homogenous from heterogeneous catalysis with electrochemical methods

In organometallic catalysis, it is often desirable to confirm the identity or composition of a molecular catalyst. An important step in identifying the “true,” or active, catalyst is distinguishing whether the materials in the electrochemical cell behaves as a homogenous or heterogeneous electrocatalyst. While there is an excellent precedent to explore heterogeneous materials in electrocatalysis because of their ease of separation from products, crucial structural, electronic, and mechanistic information is typically lost because traditional spectroscopic techniques such as NMR, IR, and electronic absorption spectroscopies are limited in the interrogation of these systems. However, there are several techniques and methods that may be used to investigate catalytic systems in an effort to distinguish homogenous from heterogeneous catalysts.102 These methods include strategic catalyst poisoning,103 X-ray photoelectron spectroscopy,104 and transmission electron microscopy.105 However, the in situ generation of new molecular species during electrocatalysis provides a unique challenge when there is a concern about distinguishing between homogeneous and heterogeneous catalysis. Distinguishing between homogeneous and heterogeneous electrocatalysts may be accomplished with electrochemical techniques such as a two-cell CV experiment, or by using an electrochemical quartz crystal microbalance (EQCM).105 In a two-cell CV experiment, homogeneous catalysis is distinguished from heterogeneous catalysis by running a series of CV scans under catalytic conditions. The first cell will contain the candidate catalyst, substrate, and electrolyte solution. When running this experiment, typical catalytic current enhancement should be observed. Following this series of scans, the working electrode is transferred to a second electrochemical cell which contains only substrate and electrolyte solution. A series of CVs identical to those run in the first cell are then run. If there is still a large current enhancement observed, this suggests the deposition of heterogeneous material behaving as the catalyst at the surface of the electrode. If the current enhancement is absent in the fresh solution, this provides support for molecular homogeneous catalysis. Another method for distinguishing heterogeneous from homogeneous catalysis is by using an electrochemical quartz crystal microbalance (EQCM).105–107 This piezoelectric gravimetric technique can detect small mass changes at the surface of the working electrode as a function of scan rate during common electrochemical experiments, such as CV. In the case of formation of a heterogeneous catalyst, a change in mass at the surface of the electrode would be detectable during a catalytic CV experiment. Conversely, for a homogenous system, there would not be a detectable change in the mass at the working electrode surface. An example from our group highlights the possible distinction of heterogeneous catalysis from homogeneous catalysis in the analysis of a well-known molecular cobaloxime complex which catalyzes hydrogen production. The deposition of heterogeneous cobalt material on the surface of the working electrode could be detected during a CV experiment with an EQCM and is distinctive from an electrochemical blank containing only electrolyte solution (See Fig. 16).

1.09.5.3

Investigating organometallic electrocatalysis using controlled potential electrolysis

To generate the product(s) associated with the current enhancement observed during electrocatalytic CV experiments, CPE experiments are commonly employed. For a detailed explanation of bulk and controlled potential electrolysis methods, Bard and Faulkner dedicate a chapter to the discussion of these techniques.36 Briefly, a CPE is typically carried out in an electrochemical cell designed specifically for the product being analyzed (See Fig. 17). Common design principles among the CPE cells include two separated compartments, with a known volume and headspace, capable of a gas-tight seal, and the ability to have headspace withdrawn after electrolysis is complete. Using the first CPE cell as an example, a typical experiment is set up where the left compartment contains the candidate electrocatalyst, the substrate, electrolyte solution, a large surface area working electrode, and reference electrode, while the right

270

Electrochemistry in Organometallic Chemistry

Fig. 16 Upper panel: cyclic voltammetry data (pink trace) and gravimetry data (black trace) showing involvement of heterogeneous, electrodeposited material under potentials at which nominally molecular catalysis is indicated by CV data alone. See reference 105 for details. Lower scheme: After the addition of acid to a (bis-difluoroboryl)cobaloxime, spectral evidence from chemical work supports the presence of the demetallated, protonated macrocycle and cobalt(hexakisacetonitrile) ([Co(NCMe)6]2+). When conducting a CV of (bis-difluoroboryl)cobaloxime in the presence of acid during an EQCM experiment, the observation of a catalytic wave is observed, as well as the deposition of cobalt metal on the surface of the working electrode. The cobalt metal is generated by reduction of [Co(NCMe)6]2+, a reaction that can be probed by independent electrochemical work with [Co(NCMe)6]2+. Modified from Sconyers, D. J.; Blakemore, J. D. Distinguishing between homogeneous and heterogeneous hydrogen-evolution catalysis with molecular cobalt complexes. Chem. Commun. 2017, 53 (53), 7286–7289.

compartment contains a sacrificial reductant (for reductive catalysis) or oxidant (oxidative catalysis), electrolyte solution, and the counter electrode. The experiment is set up to apply a bias, ideally at the determined Ecat/2 from the electrocatalytic CV experiments, for a fixed amount of time, typically ranging from 30 min, up to 24 h. Once the experiment is complete, the amount of product produced may be quantified via various detection methods. Based on the amount of charge passed during the electrolysis, and the amount of product detected, the Faradaic efficiency, turnover number, and turnover frequency can also be calculated. The Faradaic efficiency is determined by converting the charge passed, into moles of electrons passed, and then into the theoretical amount of product that should be produced. The actual yield is then divided by this theoretical yield to obtain the FE. The TON is determined by computing the moles of product per moles of catalyst. Finally, the TOF is determined by dividing the amount of product produced over unit time.

Electrochemistry in Organometallic Chemistry

271

Fig. 17 Example of controlled potential electrolysis cells that can be used to conduct WO, AO, HER, CO2R, N2RR, or EOT reactions. Cells can be custom made to specialize in gaseous product detection (left108) or ordered directly from a supplier if the goal is complete liquid or solid isolation (right, from redoxme109). The bulk electrolysis cell on the right may be available for purchase at redox.me; item: bulk electrolysis basic cell – 50 mL. https://redox.me/products/bulk-electrolysis-basiccell-50-ml. Accessed June 2021.

1.09.5.4

Product analysis in organometallic electrochemistry

To analyze the product(s) produced during organometallic electrocatalysis, quantitative methods of analysis are employed.110 Products generated in electrocatalytic reactions may be gaseous, liquid, or solid. Common laboratory techniques and instruments such as gas chromatography (GC), liquid chromatography-mass spectrometry (LC-MS), nuclear magnetic resonance (NMR), electron paramagnetic resonance (EPR), UV-Vis, and gravimetry may be used to determine the amount of product generated during electrocatalytic experiments. For gaseous product detection and quantification of common electrocatalytic products such as O2, N2, H2, and CO, gas chromatography is capable of rapid product detection and quantification when a calibration curve can be established for the particular product being examined. Likewise, when samples contain liquid or dissolved solids in solution, LC-MS techniques can be used to identify and quantify products. Furthermore, solid and liquid samples that can be dissolved in deuterated solvents may also be examined via NMR and integrated in reference to a known standard, such as trimethoxybenzene, to quantify the product produced. For paramagnetic products, EPR can be used to quantify the amount of product produced by using spin quantitation.111 Some electrolysis products, such as ammonia, can undergo a quantitative reaction that gives a product that can be quantitated colorimetrically, using UV-Vis since the molar absorptivity of the generated product is known. Lastly, for solid samples that cannot be reasonably dissolved, gravimetric methods may be used to weigh out the product produced from the CPE experiment.

1.09.6

Applications of organometallic electrocatalysis for select transformations

1.09.6.1

Electrocatalytic water oxidation

Solar fuel cells are one possible solution to the renewable energy storage challenge, and the basic design makes use of a solar photovoltaic assembly used in combination with an efficient WO electrocatalyst at the anode and a HER electrocatalyst at the cathode to generate H2.89 The allure of a clean energy future powered by sunlight and water continues to spark interest in electrocatalytic WO, which has been the topic of many reviews.112–115 As inspiration, metalloenzymes found in nature often serve as role models for the development and investigation of organometallic electrocatalysts. For instance, the oxygen-evolving complex (OEC) in photosystem II is responsible for oxidizing water to dioxygen, with release of four protons and four electrons.116 When written according to the standard of the IUPAC, the reduction potential for the interconversion of water and oxygen is +1.23 V vs SHE (see Table 1).32,117 However, if you wish to discuss the potential for the oxidation, which may be useful in computing electrochemical cell potentials, the value will be −1.23 V; this is the minimum voltage needed to drive WO. For example, using a platinized platinum electrode as a cathode, the equilibrium potential of the H2/H+ couple is set to 0 V vs SHE. At the anode, two equivalents of water is oxidized to dioxygen, four protons, and four electrons at +1.23 V vs SHE. Thus, using Eq. (2), the four electron oxidation of water to dioxygen and dihydrogen is shown to be endothermic by approximately 113 kcal/mol under aqueous conditions (see Eqs. 7–10).113 Overall : 2H2 O ! O2 + 2H2

(7)

272

Electrochemistry in Organometallic Chemistry

Table 1 Standard thermodynamic electrochemical half-reaction potentials for water in aqueous solution.

Thermodynamic half-reactions

Potential, Eo (V vs SHE)

O2 + 4H+ + 4e− . 2H2O H2O2 + 2H+ + 2e− . 2H2O HO• + H+ + e− . H2O HOO• + 3H+ + 3e− . 2H2O

+1.23 +1.76 +2.38 +1.65

Modified table from Bard, A. J.; Jordan, J.; Parsons, R. Standard Potentials in Aqueous Solutions; Marcel Dekker: New York, 1985.

Cathode : 2H + + 2e − ! H2 +

(8)

Anode : 2H2 O ! O2 + 4H + 4e   C kJ kcal ¼ + 113 ð0 − 1:23 V Þ ¼ + 475 DG ¼ − nFDE ¼ − ð4 mol e − Þ 96; 485 mol e − mol mol −

(9) (10)

The calculation shown here is readily extended to other reactions too, but with the caveat that the reversible potentials for the reactions at the cathode and the anode must be known for a specific set of conditions that satisfy the Nernst equation. Partial oxidation of water is also possible, generating products such as hydroxyl radicals, hydrogen peroxide, and peroxyl radicals, along with the appropriate number of protons and electrons. The mechanisms for formation of these partial oxidation products are also important to understand, because they may be generated as undesired side products during WO. Heterogeneous metal oxides of manganese, ruthenium, and iridium have long been recognized as electrocatalysts for WO.118–120 On the other hand, homogeneous catalysts were discovered much more recently; such homogeneous systems offer numerous advantages with regard to mechanistic studies, detection of reactive intermediates, and development of tunable catalysts. In 1982, Meyer and co-workers were the first to synthesize and characterize a molecular ruthenium electrocatalyst for wateroxidation.121 Since then, other homogeneous and heterogeneous molecular catalysts of ruthenium,122,123 iridium,124–126 and other metals,127,128 have allowed for more in-depth analysis of the steps involved in WO. For example, Brudvig, Crabtree, and co-workers,129 Macchioni and co-workers130 and several other groups have developed iridium-based catalyst precursors supported by Cp or Cp rings along with a variety of other ligands, including diimines, halides, waters, carbonyls, and phosphines (see Fig. 18 for some example complexes). Most of these complexes serve as precatalysts for water oxidation, in that initial electrochemical oxidation generates the active catalytic species that can then undergo further oxidation, triggering catalytic water oxidation. One notable example precatalyst, [Cp Ir(H2O)3]2+, undergoes oxidation to form a heterogeneous “blue layer” on the electrode surface that is readily observed by a variety of techniques, including cyclic voltammetry.131 In other cases, catalyst activation leads to formation of homogeneous species that can catalyze water oxidation in solution (i.e., without formation of solids or particulate species). Distinguishing between homogeneous and heterogeneous catalysis, a recognized challenge in organometallic chemistry, is made especially challenging under these conditions for water oxidation catalysis driven by organometallic precatalysts. The electrochemical quartz crystal microbalance has been shown to be a useful tool for observing formation of insoluble heterogeneous

Fig. 18 Select iridium(III) organometallic precatalysts for electrochemical WO. From Blakemore, J. D.; Schley, N. D.; Balcells, D.; Hull, J. F.; Olack, G. W.; Incarvito, C. D.; Eisenstein, O.; Brudvig, G. W.; Crabtree, R. H. Half-Sandwich Iridium Complexes for Homogeneous Water-Oxidation Catalysis. J. Am. Chem. Soc. 2010, 132 (45), 16017–16029 with permission.

Electrochemistry in Organometallic Chemistry

273

species, both for oxidative reactions like water oxidation driven by organometallic precursors132 as well as reductive reactions like hydrogen evolution and metal electrodeposition.104,106,107 Along a similar line, dynamic light scattering is useful for formation of particulates or particles in solution. Application of specialized techniques like these, in concert with detailed electroanalytical and mechanistic/kinetic/chemical work, can strongly inform studies of water oxidation catalysis. Approaching molecular electrocatalytic systems from multiple viewpoints and with multiple chemical and electrochemical techniques can thus be an appealing strategy, especially when dealing with challenging reactions like water oxidation.

1.09.6.2

Electrocatalytic ammonia oxidation

Electrocatalytic AO and WO are comparable in many ways; both reactions involve the removal of several electrons and protons, the formation of a new bond between heteroatoms, and the evolution of a gaseous product (N2 or O2). Likewise, electrocatalytic AO represents an alternative fuel cell design where AO would be carried out at an anode and proton reduction at a cathode to generate energy. However, unlike WO, electrocatalytic AO is still an emerging field and has been the topic of a recent review by Wu and co-workers133 and perspective by Bullock and co-workers.134 Complete ammonia oxidation is best summarized as a thermodynamic electrochemical half reaction where ammonia is converted into dinitrogen, six protons, and six electrons at −0.057 V vs SCE (see Table 2). Partial ammonia oxidation is also possible, generating other products such as hydrazine and hydroxylamine along with the relevant number of protons and electrons, on route to complete ammonia oxidation. Heterogeneous materials have been employed in electrocatalytic AO, but these instances rely heavily on precious metals and thus proposed mechanisms are not secure. The first report of a molecular electrocatalyst for AO was reported by Smith, Hamann, and co-workers, where they used a derivative of a Ru-based WO catalyst bearing an electron-donating Me2Nbpy ligand.135 Subsequently, other electrocatalysts for water oxidation have been repurposed for ammonia oxidation. Recent work from the Peters group makes use of a previously reported alkane oxidation catalyst.136 The labile acetonitrile (MeCN) ligands of the starting complex are readily exchanged in the presence of ammonia, shown here in the solid-state structure of [(bpyPy2Me)Fe(MeCN)(NH3)](OTf )2 (see Fig. 19). Compared to [(TPA)Fe(MeCN)2](OTf )2 (TPA is tris(2-pyridylmethyl)amine),99 a catalytic CV experiment shows a dramatic increase in a catalytic current for this newly reported complex and FOWA was performed to determine the observed rate of AO. The CPE cell was loaded with 0.05 mM [Fe] catalyst, 20 mM of NH3, and 50 mM NH4OTf in MeCN as the supporting electrolyte. After applying a bias of 0.85 V vs Fc+/0 for 24 h, product analysis revealed the formation of N2 at the working electrode and H2 at the counter electrode in all instances, as well as the highest TON to date for a molecular AO electrocatalyst. Future work in the field of electrocatalytic AO requires probing the rational design, kinetics, and thermodynamics of molecular complexes.

1.09.6.3

Electrocatalytic proton reduction

In the electrocatalytic HER, two protons and two electrons are coupled with the help of a catalyst to generate dihydrogen. This is simpler than the multielectron N2 and O2 reactions, electrocatalytic HER has been the topic of significant work spanning decades. Electrocatalytic HER continues to draw major interest in the solar fuels community because understanding the elementary steps involved in this simple energy storage reaction could help in the rational design of future electrocatalysts.137,138 Complete proton reduction to dihydrogen is represented by the thermodynamic electrochemical half reaction where two protons and two electrons are coupled together to generate dihydrogen at 0.00 V vs SHE (see Table 3).32 Gray, Winkler, Brunschwig and co-workers have discussed possible mechanisms for dihydrogen evolution,137 as have other researchers in this vibrant area of electrocatalysis research. HER catalysis typically involves either a homolytic or heterolytic pathway to the production of dihydrogen. The heterolytic pathway involves a proton and a hydride reacting to generate dihydrogen, while the homolytic pathway involves homolytic cleavage of two hydride species which undergo recombination to produce dihydrogen. Heterogeneous noble metals, such as platinum, are prominent electrocatalysts for the reduction of protons to dihydrogen. Table 2 Standard thermodynamic electrochemical half-reaction potentials for dinitrogen in aqueous solution.

Thermodynamic half-reactions24

Potential Eo (V vs SHE)

N2 + 6H+ + 6e− . 2NH3 N2 + 8H+ + 6e− . 2NH+4 N2 + 4H2O + 4e− . N2H4 + 4OH− N2 + 5H+ + 4e− . N2H+5 3N2 + 2e− . 2N−3

−0.06 +0.28 −1.16 −0.23 −3.40

Modified table from Bard, A. J.; Jordan, J.; Parsons, R. Standard Potentials in Aqueous Solutions; Marcel Dekker: New York, 1985.

274

Electrochemistry in Organometallic Chemistry

Fig. 19 Solid-state structure of [(bpyPy2Me)Fe(MeCN)(NH3)](OTf )2 (left). Thermal ellipsoids shown at 50% probability. Triflate and hydrogen atoms are omitted for clarity. CV of MeCN solutions containing 400 equiv. of NH3, 0.05 M NH4OTf, 0.5 M of [(bpyPy2Me)Fe(MeCN)2](OTf )2 or [(TPA)Fe(MeCN)2](OTf )2. From Zott, M. D.; Peters, J. C. Enhanced Ammonia Oxidation Catalysis by a Low-Spin Iron Complex Featuring Cis Coordination Sites. J. Am. Chem. Soc. 2021, 143 (20), 7612–7616 with permission.

Table 3 Standard thermodynamic electrochemical half-reaction potentials for protons in aqueous solution.

Thermodynamic half-reaction

Potential Eo (V vs SHE)

2H+ + 2e− . H2

0.00

Modified table from Bard, A. J.; Jordan, J.; Parsons, R. Standard Potentials in Aqueous Solutions; Marcel Dekker: New York, 1985.

139 Other heterogeneous materials, such as metal nanoparticles, have also shown promise for the electrocatalytic production of dihydrogen.140 Molecular cobaloxime complexes have also been shown to be highly active for the production of dihydrogen.141 While these systems are competent and efficient HER electrocatalysts, they may generate Co nanoparticles in the presence of acid which generates ambiguity in the identity of the true electrocatalyst.105 Thus, well-defined homogeneous model systems that show activity toward electrocatalytic HER have been desirable. While many redox-active systems are capable of electrocatalytic HER, the homogeneous Rh-based electrocatalyst, [Cp RhCl H ( bpy)]+, developed by Kölle and Grätzel is one of the most well-known.142 CV of [Cp RhCl(Hbpy)]+ reveals a quasi-reversible 2e− redox couple centered at approximately −1.21 V vs Fc+/0. Addition of two equivalents of acid to Cp Rh(bpy) results in the quantitative production of dihydrogen, and further experiments confirm its catalytic ability. This led to a systematic examination of a series of [Cp RhCl(Rbpy)]+ (Rbpy is 4,40 -disubstituted-2,20 -bipyridyl; R ¼ H, tBu, and CF3) complexes by Blakemore and co-workers in an effort to determine if electrocatalytic dihydrogen evolution could be modulated using substituents on the bpy ligand.73 Initial catalytic CV experiments for these complexes were discussed in Section 1.09.5.2.1. These experiments assisted in the determination of the parameters needed to carry out a series of CPE experiments. To systematically compare the catalytic ability of each complex, CPE experiment were conducted with identical conditions with 1 mM of [Rh] catalyst and 10 mM anilinium triflate in the WE portion of the cell, ferrocene was added to the CE portion of the cell as a sacrificial reductant, 0.1 M TBAPF6 in MeCN was used as the supporting electrolyte, and a bias of −1.36 V vs Fc+/0 was applied for 90 min. Over the course of the experiment, steady current flow is observed for each complex over time (see Fig. 20). Notably, each of the complexes has significantly more current flowing than the blank and the amount of current flowing is different based on the identity of the complex. When the experiment is first started, the first few seconds appears to show rapid consumption of acid, but this is actually the initial charging of the WE, and following this period, steady current flows over time. Upon completion of the experiment, the presence of dihydrogen was confirmed in all instances by GC. When the identity of the substituents was –tBu, –H, or –CF3, the complexes produced 2.6 mL, 3.5 mL, and 1.5 mL of dihydrogen, with corresponding TONs of 3.4, 4.4, and 2.7, respectively.

Electrochemistry in Organometallic Chemistry

275

Fig. 20 [Cp RhCl(Rbpy)]+ complexes with electron donating and withdrawing substituents (left). Comparison of CPE steady-state current as a function of time with an applied bias of −1.36 V vs Fc+/0 with 10 mM anilinium triflate as the proton source (right). Modified from Henke, W. C.; Lionetti, D.; Moore, W. N. G.; Hopkins, J. A.; Day, V. W.; Blakemore, J. D. Ligand Substituents Govern the Efficiency and Mechanistic Path of Hydrogen Production with [Cp Rh] Catalysts. ChemSusChem 2017, 10 (22), 4589–4598.

As an aside, keeping the bias at the same potential across the series of complexes results in what could be perceived as an “unfair” although uniformly direct comparison of these catalysts’ abilities to produce dihydrogen. This is because from the perspective of the Nernst equation, when the potential is moved from Ecat/2 of a given individual complex, the amount of reduced species present at the electrode at any one point in time is changed. In other words, catalysis for each of these systems will be optimal when polarizing at Ecat/2, and thus comparing the kinetic performance of the catalysts at a common potential value that doesn’t align with the value of Ecat/2 for all the catalysts introduces an additional concentration/potential dependence to this study. On the other hand, a common potential polarization for all the comparisons made here does satisfy a practical aspect of testing and comparing these model catalysts under comparable conditions. Regardless of this potential ambiguity, however, based on the amount of charge transferred and the amount of dihydrogen detected, all three complexes were confirmed (in this study) to be electrocatalysts with FE exceeding 90% in all three systems. This study encapsulates the challenges and opportunities in comparing across a family of catalysts that may have unique redox chemistries.

1.09.6.4

Electrocatalytic carbon dioxide reduction

Electrocatalytic CO2 reduction continues to draw attention from the organometallic chemistry community because organometallic compounds are uniquely well-suited to understanding the bonding and activation of carbon-containing small molecules.143 Use of CO2 as input for preparation of useful chemicals could also decrease the environmental impact of chemical industry, motivating work in this area further. CO2 has many potential reduction pathways and products, with one of the most common targets being carbon monoxide (CO) because of its use as an industrial synthon. More elaborate reactions could be used to produce carbon-based fuels or other chemicals. The thermodynamic electrochemical half-reaction for some key CO2 reduction products involving two, four, six, or eight electron pathways are shown below (see Table 4).32,144 Heterogeneous electrode materials such as carbon, copper, and gold are capable of reducing CO2 directly, but with various reaction pathways and significant overpotentials.145 The use of these materials precludes the use of spectroscopic methods, making the development of future heterogeneous materials more difficult. Many homogeneous molecular complexes bearing first- and third-row transition metals have been synthesized with the goal of using them as CO2 reduction catalysts. A notable example from Savéant and co-workers uses a molecular catalyst, tetraphenylporphyrin Fe(III) chloride (TPPFe(III)Cl), which undergoes three sequential one electron reductions to generate the active electrocatalyst for CO2 reduction (see Fig. 21).146 In the presence of trifluoroacetic acid and CO2, the reversible behavior in the CV of the Fe(TPP) complex ceases and a substantial catalytic wave is observed at around −1.5 V vs SCE. Preparative scale CPE results in the selective production of CO at 96% FE. However, the tricarbonyl complexes of the group 7 metals Re and Mn are perhaps the most famous CO2 reduction electrocatalysts. The initial design of this class of homogeneous molecular CO2 electrocatalysts began in the 1980s with the initial report of fac-Re(CO)3Cl(Hbpy) by Lehn and co-workers147 and Meyer and co-workers.148 Approximately 30 years later, Deronzier, Chardon-Noblat, and co-workers introduced the manganese analog of the complex, fac-Mn(CO)3Br(Hbpy) as a significantly more Earth-abundant electrocatalyst for CO2 reduction.68 Both fac-Re(CO)3Cl(Hbpy) and fac-Mn(CO)3Br(Hbpy) are comparable and competent catalysts for the generation of CO from CO2. Recent work in this area has focused on tuning the electron density at the [Re] and [Mn] metal centers by using 4,40 -disubstituted-bpy (Rbpy) ligands.87,101,149

276

Electrochemistry in Organometallic Chemistry

Table 4 Standard thermodynamic electrochemical half-reaction potentials for CO2 in aqueous solution.

Thermodynamic half-reactions

Potential Eo (V vs SHE)

CO2 + 2H+ + 2e− . CO + H2O CO2 + 2H+ + 2e− . HCO2H CO2 + H2O + 2e− . HCO−2 + OH− CO2 + 2H2O + 2e− . CO + 2OH− CO2 + 4H+ + 4e− . CH2O + H2O CO2 + 3H2O + 4e− . CH2O + 4OH− CO2 + 6H+ + 6e− . CH3OH + H2O CO2 + 8H+ + 8e− . CH4 + 2H2O CO2 + 5H2O + 6e− . CH3OH + 6OH− CO2 + 6H2O + 8e− . CH4 + 8OH− 2CO2 + 2H+ + 2e− . C2H2O4 2CO2 + 2e− . C2O2− 4 2CO2 + 12H+ + 12e− . C2H4 + 4H2O 2CO2 + 12H+ + 12e− . C2H6O + 3H2O CO2 + 4H+ + 4e− . C + 2H2O CO2 + 2H2O + 4e− . C + 4OH−

−0.106 −0.199 −1.078 −0.934 −0.070 −0.898 +0.016 +0.169 −0.812 −0.569 −0.475 −0.590 +0.064 +0.084 +0.206 −0.627

Modified table from Bard, A. J.; Jordan, J.; Parsons, R. Standard Potentials in Aqueous Solutions; Marcel Dekker: New York, 1985.

Fig. 21 CV of TPPFe(III)Cl at 100 mV s−1 using 0.1 M TEAClO4 in dimethylformamide as the supporting electrolyte and a glassy carbon working electrode. The sequential reductions of TPPFe(III)Cl at these negative potentials suggests that it was a prime candidate for electrocatalytic CO2 reduction. From Bhugun, I.; Lexa, D.; Savéant, J.-M. Ultraefficient selective homogeneous catalysis of the electrochemical reduction of carbon dioxide by an iron(0) porphyrin associated with a weak Broensted acid cocatalyst. J. Am. Chem. Soc. 1996, 118 (7), 1769–1776 with permission.

The high cost of Re has resulted in an influx of reports on fac-Mn(CO)3Br(Rbpy) complexes. Some fac-Mn(CO)3Br(Rbpy) derivatives have been used to address potential shortcomings of Mn(I) complexes, including the known visible-light photosensitivity and redox-induced dimerization of fac-Mn(CO)3Br(Rbpy) to form Mn2(CO)6(Rbpy)2.150,151 Kubiak and co-workers demonstrated that the dimerization of fac-Mn(CO)3Br(Rbpy) complexes can be impeded by utilizing bulky bipyridine ligands bearing mesityl groups in the 6 and 60 positions of the bpy ligand (mesbpy). CV of fac-Mn(CO)3(NCMe)(mesbpy) reveals a quasi-reversible 2e− redox couple centered at −1.55 V vs Fc+/0 (see Fig. 22).152 In a catalytic CV experiment, fac-Mn(CO)3 (NCMe)(mesbpy) is reduced in the presence of methanol and CO2. Interestingly, the onset of the catalytic wave does not occur at the initial 2e− reduction, but rather, there is a kinetic potential shift that indicate the binding of CO2, before being reduced again at −2.0 V to induce catalysis. Preparative CPE experiments reveal fac-Mn(CO)3(NCMe)(mesbpy) is a competent CO2 reduction electrocatalyst and selectively produces CO at 98% FE. Future work in the area of electrocatalytic CO2 reduction will likely focus on kinetic, mechanistic, and thermodynamic studies in areas that extend beyond the common two-electron reduction products.

Electrochemistry in Organometallic Chemistry

277

Fig. 22 Catalytic CV experiment demonstrating that catalysis only occurs when fac-Mn(CO)3(NCMe)(mesbpy) is in the presence of CO2 and MeOH (left). A proposed catalytic cycle for the electrocatalytic generation of CO using Mn(CO)3(NCMe)(mesbpy) (right). From Sampson, M. D.; Nguyen, A. D.; Grice, K. A.; Moore, C. E.; Rheingold, A. L.; Kubiak, C. P. Manganese Catalysts with Bulky Bipyridine Ligands for the Electrocatalytic Reduction of Carbon Dioxide: Eliminating Dimerization and Altering Catalysis. J. Am. Chem. Soc. 2014, 136 (14), 5460–5471 with permission.

1.09.6.5

Electrocatalytic dinitrogen reduction

Functionalizing dinitrogen is a particularly challenging process, and this is manifested in its physicochemical characteristics.153 However, since dinitrogen comprises more than 79% of the Earth’s atmosphere, chemical or electrochemical generation of reduced nitrogen products, such as ammonia, could be a useful way to utilize this abundant resource. On an industrial scale, nitrogen fixation is accomplished using the Haber-Bosch process where dihydrogen (produced from the steam reformation of methane), dinitrogen (obtained from air), and a doped heterogeneous Fe-based catalyst are used to produce ammonia.154 The Haber-Bosch process itself is quite efficient, converting >97% of the chemical inputs into ammonia, a fact made more incredible when considering that this reaction accounts for nearly half of the Earth’s total ammonia production including both natural and artificial reactivity. However, the Haber-Bosch process as currently performed is the source of copious amounts of CO2, and requires reaction temperatures exceeding 400  C and pressures surpassing 200 atm. As a result, the redox process of reducing dinitrogen to ammonia has attracted increased interest from the electrocatalysis community; from this perspective, ammonia could either be used in fuel cells, producing only dinitrogen and water in an ideal scenario, or in its traditional role as a crop fertilizer. In nature, nitrogen fixation occurs by using the enzyme nitrogenase to reduce dinitrogen at room temperature and ambient pressure.155 Toward this goal, the thermodynamic electrochemical half-reactions for some key dinitrogen reduction products involving, two, four, or six electrons are shown above (see Table 2 in AO) (Fig. 23).32 While the electrocatalytic transformations (WO, HER, and CO2R) discussed so far have many examples, electrocatalytic N2RR with homogeneous catalysts is rare.156 The first report of a true electrocatalytic N2RR came from the Peters group in 2016 using [(PB3)Fe]+ as a precatalyst and HBArF as an acid source.157 Catalytic CV experiments revealed the onset of a catalytic wave at the [(PB3)Fe]+/[(PB3)Fe(N2)] couple at −1.5 V vs Fc+/0, but greater enhancement was observed at the [(PB3)Fe(N2)]0/−1 couple at −2.2 V vs Fc+/0. This potential is well within the range to observe HER as well as N2RR. CPE of this complex revealed the presence of ammonia (2.3 equiv.) (as well as substantial amounts of H2), confirming the catalytic ability of the complex. More recently, better kinetic control was achieved to improve the selectivity and turnover of this reaction by using a metallocene molecular mediator to facilitate the reduction of dinitrogen.158,159 Improvement was achieved through the use of Cp 2Co as a mediator, resulting in a catalytic wave with onset at  −2.0 V vs Fc+/0. CPE of [(PB3)Fe]+, with [Ph2NH2]+ and a cocatalytic amount of Cp2Co, at an applied bias of −2.1 V vs Fc+/0 resulted in 4 equiv. of ammonia. Future work in this area depends on the synthesis of new heterogeneous or homogeneous electrocatalysts capable of N2RR. This will be accomplished by developing systems and conditions that can operate negative enough to engender N2RR, but at the same time disfavor the kinetics that prefers the HER (Fig. 24).

Fig. 23 Reduction of dinitrogen using protons and electrons in combination with an electrocatalyst to generate ammonia. Modified table from Bard, A. J.; Jordan, J.; Parsons, R. Standard potentials in aqueous solutions; Marcel Dekker: New York, 1985.

278

Electrochemistry in Organometallic Chemistry

Fig. 24 First molecular electrocatalyst for N2RR in the presence of a cocatalytic amount of [Cp 2Co]+. From Chalkley, M. J.; Drover, M. W.; Peters, J. C., Catalytic N2-to-NH3 (or -N2H4) Conversion by Well-Defined Molecular Coordination Complexes. Chem. Rev. 2020, 120 (12), 5582–5636 with permission.

1.09.6.6

Electrocatalytic organic transformations

Organometallic electrochemistry is typically associated with redox-induced small molecule activation where electrocatalysts are used to access and store energy. However, these techniques also are readily extended to electrocatalytic organic (electroorganic) transformations, which are rapidly gaining popularity.160 Fortunately, organic chemists are beginning to think of electrochemistry as an essential method for driving synthetic processes. In this field, organometallic electrocatalysts typically behave as redox mediators, shuttling electrons from the surface of the working electrode to an organic substrate on the way to product generation.151,161 In some instance, these electrocatalysts may also engender regio-, chemo-, and stereoselectivity in the products (see Fig. 25).162 In these reactions organic substrates are oxidized or reduced at a controlled potential, resulting in the generation of intermediate radical species that undergo further chemical reactivity before generating the product(s). To this end, electroorganic syntheses mediated by organometallic complexes have attracted the interest of commodity chemical and pharmaceutical industries because of the tunability, scalability, and potential energy control provided by electrochemical techniques. Indeed, selective manipulation and installation of functional groups on organic substrates with electrochemical methods offers a distinct new approach that is likely to continue surging in activity. The earliest preparative scale electrochemical experiment was performed by Faraday in 1834, where an acetic acid solution was placed under anodic oxidative bias to generate ethane.163 This experiment went on to inspire the Kolbe electrolysis in 1847, where abundant carboxylic acids were oxidized and used to produce alkyl radicals.164 Around the same time, the first reductive preparative scale electroorganic experiment was also reported; this involved the reductive dehalogenation of trichlormethane sulfonic acid to methanesulfonic acid.165 In the last 20 years, electrochemical methods have become more accessible to synthetic organic chemists due to the widespread availability of potentiostats and advancements in electrochemical techniques and methods, and this has resulted in an uptick in catalytic electroorganic synthesis literature reports.166 Common functional groups such as aldehydes, ketones, esters, and olefins have been studied under electrocatalytic oxidative and reductive conditions with success. Hydrogenation of olefins and alkynes has also been studied, taking advantage of reactive transition metal hydrides typically implicated in HER electrocatalysis.167 A common theme in organic electrocatalysis is the use of a redox mediator to drive a chemical transformation. In an example of organometallic electroorganic synthesis, Xu and co-workers have taken advantage of the reliable one-electron redox chemistry of ferrocene.168 In their work, ferrocene is a mediator in the electrocatalytic C–H and N–H functionalization of functionalized (aza) indoles. A catalytic CV experiment shows that the addition of the model urea-based substrate in the presence of ferrocene does not result in any significant change to the appearance of the CV which displays the reversible oxidation of ferrocene to ferrocenium (see

Fig. 25 (A) a molecular mediator (CAT), in the context of oxidation, donating an electron to the working electrode, and later being reduced by the substrate (SUB) in route to product generation. (B) a molecular mediator, in the context of oxidation, donating an electron to the working electrode, stabilizing the substrate to form a catalyst-substrate complex, and then generating an intermediate on route to product generation. From Siu, J. C.; Fu, N.; Lin, S., Catalyzing Electrosynthesis: A Homogeneous Electrocatalytic Approach to Reaction Discovery. Acc. Chem. Res. 2020, 53 (3), 547–560 with permission.

Electrochemistry in Organometallic Chemistry

279

Fig. 26 A proposed mechanism for the catalytic electrosynthesis of highly substituted indoles (right). From Yan, M.; Kawamata, Y.; Baran, P. S., Synthetic Organic Electrochemical Methods Since 2000: On the Verge of a Renaissance. Chem. Rev. 2017, 117 (21), 13230–13319 with permission.

Fig. 26). However, when sodium methoxide is added, a catalytic wave is observed and the return reduction has ceased, suggesting that the ferrocenium is consumed by the urea-based substrate. A possible mechanism involves a pre-equilibrium with deprotonation of the urea-based complex, resulting in a nitrogen-centered anion which then undergoes oxidation by intermolecular electron transfer with the electrogenerated ferrocenium. The oxidation generates a nitrogen centered radical that is capable of intramolecular cyclization with the neighboring alkyne which ultimately results in the formation of highly substituted indoles, while also regenerating ferrocene. Electrochemical methods for this redox-induced synthetic organic chemistry are advantageous because waste byproducts associated with the use of stoichiometric oxidants and reductants are eliminated. Future work in this area needs to focus on systematic studies of appropriate redox mediators capable of driving more complicated organic transformations. Important reactions, such as those involving the activation of aliphatic CdH bonds near amine and amide functionalities, is currently an active area in electroorganic synthesis.169

1.09.6.7

Analyses for benchmarking electrocatalysts

The significant increase in the production of new electrocatalysts over the years has contributed to the need for methods that can reliably benchmark in order to determine what is a “good” electrocatalyst. To this end, the derived Butler-Volmer36 and Tafel170 equations are two common analyses for benchmarking electrocatalysts, especially in the realm of heterogeneous catalysis. These methods are typically used to benchmark WO, HER, and CO2RR catalysts, but can be performed for AO, N2RR, and EOT reactions. Demonstrating high rates of activity and low overpotentials are the desired outcomes of catalyst benchmarking efforts. The main assessment methods are Butler-Volmer and Tafel analyses, and the reader may consult resources that give the full details.36,171 Briefly, the Butler-Volmer equation describes the relationship between the current and potential between the surface of the electrode and the bulk sample for an elementary redox reaction. (At high overpotentials, the Butler-Volmer equation simplifies to the Tafel equation.) The resulting Butler-Volmer and Tafel analyses provide kinetic and thermodynamic insight into systems by relating the rate of electrocatalysis to the overpotential.

280

Electrochemistry in Organometallic Chemistry

Table 5 Thermodynamic equilibrium potentials for the reduction of protons, oxygen, and CO2 in MeCN and DMF.166,167 Half-reactions

2H+ + 2e− Ð H2 O2 + 4H+ + 4e− Ð 2H2O CO2 + 2H+ + 2e− Ð CO + H2O

Potential, Eo (V vs Fc+/0) MeCN

Potential, Eo (V vs Fc+/0) DMF

−0.028 +1.29 −0.12

−0.662 +0.60 −0.73

In the realm of molecular electrocatalysis, alternative methods for benchmarking catalysts and understanding their kinetic behaviors are typically utilized. One such method is Foot-of-the-Wave (FOWA) analysis, a method developed by Savéant, Costentin, and co-workers.172 Summarizing this method, FOWA examines the “foot” of catalytic (and therefore irreversible) waves at points early in the onset of the catalytic current, in order to extract the instantaneous reaction rate from the voltammetric data.173 One assumption that underlies typical FOWA is that the key redox-active intermediate involved in the rate-determining step of catalysis behaves in a Nernstian fashion, and thus its concentration can be considered directly controlled at the electrode surface by the applied potential. FOWA is distinguished as a very helpful method for extraction of rate information at short times, when catalyst deactivation, substrate consumption, and/or product inhibition are avoided. Notably, when voltammetric data is collected in such a way as to avoid these issues, other analyses based on Zone Diagrams and mathematical treatments are possible. The reader is encouraged to consult resources that give the full details on these methods.29 Significant work has also been accomplished in the last few years to better understand equilibrium potentials (EoX) of feedstocks capable of interconverting chemical and electrical energy in non-aqueous, organic solvents. These measurements establish a metric for evaluating and comparing different electrocatalysts for a given process in non-aqueous solvents. Work by Roberts and Bullock has provided an accurate determination of the H+/H2 equilibrium potential for various acids in MeCN (see Table 5).91 These experiments use a Pt electrode, which is an excellent catalyst for reversible hydrogen reduction, even in organic solvents. In related work, Mayer and Helm have also established EoX values (referenced to Fc+/0 in acetonitrile and DMF) for electrocatalytic oxygen reduction and carbon dioxide reduction.174 Determining these values in non-aqueous solvents has allowed the determination of accurate potentials for catalytic responses during electrocatalysis.

1.09.7

Conclusion

The widespread availability of sensitive and robust potentiostats has led to substantial growth in the electrochemical investigation of redox processes in diverse compounds that are of interest to organometallic chemists. Modern electrochemical materials, methods and techniques have assisted in interpreting chemical reactivity that can be promoted at electrode surfaces. Cyclic voltammetry and controlled potential electrolysis experiments have provided important insights, and continue to be workhorse methods in organometallic electrochemistry. Voltammetry experiments can reveal the potentials of redox events whose nature is then probed with stoichiometric redox reagents in chemical syntheses. The thermodynamic, kinetic, and mechanistic data provided by cyclic voltammetric experiments has informed design principles that can be used to generate organometallic complexes and catalysts with tailored properties. Further, the shapes of redox waves in cyclic voltammetry responses can provide crucial information about the nature of electron transfer events and chemical reactions in solution, offering insights difficult to obtain by other techniques. This is particularly important in the field of electrocatalysis, in which efficiency and selectivity of a given catalyst can be investigated by coupling electrolysis methods to product analysis. Considering all the useful opportunities afforded by electrochemical methods, the future of organometallic electrochemistry is bright.

Acknowledgments The authors thank the US National Science Foundation (OIA-1833087), the US Department of Energy, Office of Science, Office of Basic Energy Sciences, Early Career Research Program (DE-SC0019169), and the Kansas Academy of Science for funding different aspects of our group’s work in redox chemistry and catalysis.

References 1. 2. 3. 4.

Kealy, T.; Pauson, P. A. Nature 1951, 168, 1039–1040. Wilkinson, G.; Rosenblum, M.; Whiting, M. C.; Woodward, R. B. J. Am. Chem. Soc. 1952, 74, 2125–2126. Kuwana, T.; Bublitz, D. E.; Hoh, G. Chem. Ind. (Lond.) 1959, 20, 635–636. Kuwana, T.; Bublitz, D. E.; Hoh, G. J. Am. Chem. Soc. 1960, 82, 5811–5817.

Electrochemistry in Organometallic Chemistry 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 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. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. 61.

281

Astruc, D. Electron Transfer and Radical Processes in Transition-Metal Chemistry; VCH Publishers: New York, 1995. Geiger, W. E. Organometallics 2007, 26, 5738–5765. Werlé, C.; Meyer, K. Organometallics 2019, 38, 1181–1185. Elgrishi, N.; Rountree, K. J.; McCarthy, B. D.; Rountree, E. S.; Eisenhart, T. T.; Dempsey, J. L. J. Chem. Educ. 2018, 95 (2), 197–206. Nernst, W. Z. Phys. Chem. NF. 1888, 2U (1), 613–637. Connelly, N. G.; Geiger, W. E. Chem. Rev. 1996, 6 (2), 877–910. Savéant, J. M.; Su, K. B. J. Electroanal. Chem. 1984, 171 (1), 341–349. (a) Costentin, C.; Drouet, S.; Robert, M.; Saveant, J. M. J. Am. Chem. Soc. 2012, 134, 11235–11242; (b) Costentin, C.; Saveant, J. M. ChemElectroChem 2014, 1, 1226–1236; (c) Costentin, C.; Drouet, S.; Robert, M.; Saveant, J. M. Science 2012, 338, 90–94. Appel, A. M.; Helm, M. L. ACS Catal. 2014, 4 (2), 630–633. Stock, J. T. Electrochemistry in Retrospect. In Electrochemistry, Past and Present, 390; American Chemical Society, 1989; pp 1–17. Mann, C. K.; Barnes, K. K. Electrochemical Reactions in Nonaqueous Systems; Marcel Dekker: New York, 1970; pp 403–418. Morris, M. D. In Electroanalytical Chemistry; Bard, A. J., Ed.; 7; Marcel Dekker: New York, 1974; pp 80–121. Kauffman, G. B. J. Chem. Educ. 1983, 60 (3), 185–186. Page, J. A.; Wilkinson, G. J. Am. Chem. Soc. 1952, 74 (23), 6149–6150. (a) Dessy, R. E.; Kitching, W.; Chivers, T. J. Am. Chem. Soc. 1966, 88 (3), 453–459; (b) Dessy, R. E.; Kitching, W.; Psarras, T.; Salinger, R.; Chen, A.; Chivers, T. J. Am. Chem. Soc. 1966, 88 (3), 460–467; (c) Dessy, R. E.; Chivers, T.; Kitching, W. J. Am. Chem. Soc. 1966, 88 (3), 467–470; (d) Dessy, R. E.; Stary, F. E.; King, R. B.; Waldrop, M. J. Am. Chem. Soc. 1966, 88 (3), 471–476; (e) Dessy, R. E.; King, R. B.; Waldrop, M. J. Am. Chem. Soc. 1966, 88 (22), 5112–5117; (f ) Dessy, R. E.; Weissman, P. M.; Pohl, R. L. J. Am. Chem. Soc. 1966, 88 (22), 5117–5121. Dessy, R. E.; Pohl, R. L.; King, R. B., J. Am. Chem. Soc. 1966, 88 (22), 5121–5124; (g) Dessy, R. E.; Weissman, P. M. J. Am. Chem. Soc. 1966, 88 (22), 5124–5129; (h) Dessy, R. E.; Weissman, P. M. J. Am. Chem. Soc. 1966, 88 (22), 5129–5131; (i) Psarras, T.; Dessy, R. E. J. Am. Chem. Soc. 1966, 88 (22), 5132–5135; (j) Dessy, R. E.; Pohl, R. L. J. Am. Chem. Soc. 1968, 90 (8), 1995–2001; (k) Dessy, R. E.; Kornmann, R. L.; Smith, C.; Haytor, R. J. Am. Chem. Soc. 1968, 90 (8), 2001–2004; (l) Dessy, R. E.; Pohl, R. L. J. Am. Chem. Soc. 1968, 90 (8), 2005–2008; (m) Dessy, R. E.; Wieczorek, L. J. Am. Chem. Soc. 1969, 91 (18), 4963–4974; (n) Dessy, R. E.; Charkoudian, J. C.; Abeles, T. P.; Rheingold, A. L. J. Am. Chem. Soc. 1970, 92 (13), 3947–3956; (o) Dessy, R. E.; Charkoudian, J. C.; Rheingold, A. L. J. Am. Chem. Soc. 1972, 94 (3), 738–745; (p) Dessy, R. E.; Rheingold, A. L.; Howard, G. D. J. Am. Chem. Soc. 1972, 94 (3), 746–752; (q) Dessy, R. E.; Bares, L. A. Acc. Chem. Res. 1972, 5 (12), 415–421. (a) Huebert, B. J.; Smith, D. E. J. Electroanal. Chem. 1971, 31 (2), 333–348. For later work on the reduction of cyclooctatetraene see: Petersen, R. A.; Evans, D. H., J. Electroanal. Chem. 1987, 222 (1), 129–150; (b) Smith, D. E. In Electroanalytical Chemistry; Bard, A. J., Ed.; Marcel Dekker Inc.: New York, 1966; vol. 1 pp 102–110. (a) Gutmann, V.; Schöber, G. Monatsh. Chem. Verw. Teile Anderer Wiss. 1957, 88 (2), 206–215; (b) Vlcek, A. A. Collect. Czechoslov. Chem. Commun. 1965, 30, 952–960; (c) Mašek, J. Collect. Czechoslov. Chem. Commun. 1965, 30, 4117–4126; (d) Denisovich, L. I.; Gubin, S. P.; Chapovski Yu, A.; Ustynok, N. A. Bull. Acad. Sci. USSR, Div. Chem. Sci. 1968, 891; (e) Piazza, G.; Paliani, G. Z. Phys. Chem. 1970, 71, 91–101; (f ) Mann, C. K.; Barnes, K. K. Electrochemical Reactions in Nonaqueous Systems; Marcel Dekker: New York, 1970; pp 403–418. Adams, R. N. Electrochemistry at Solid Electrodes; Marcel Dekker: New York, 1969. Booman, G. L. Anal. Chem. 1957, 29 (2), 213–218. Smith, D. E. In Electroanalytical Chemistry; Bard, A. J., Ed.; Marcel Dekker Inc.: New York, 1966; vol. 1; pp 102–110. DeFord, D. D. American Chemical Society 133rd Meeting, San Francisco, April; . Smith, D. E. Anal. Chem. 1963, 35 (12), 1811–1820. Schwarz, W. M.; Shain, I. Anal. Chem. 1963, 35 (12), 1770–1778. Nicholson, R. S.; Shain, I. Anal. Chem. 1964, 36 (4), 706–723. Savéant, J.-M.; Costentin, C. Elements of Molecular and Biomolecular Electrochemistry: An Electrochemical Approach to Electron Transfer Chemistry, 2nd ed; John Wiley & Sons: Hoboken, NJ, 2019. Pourbaix, M. Atlas of Electrochemical Equilibria in Aqueous Solutions, 2nd ed.; National Association of Corrosion Engineers: Houston, 1974. Latimer, W. M. Oxidation Potentials, 2nd ed.; Prentice-Hall: New York, 1952. Bard, A. J.; Jordan, J.; Parsons, R. Standard Potentials in Aqueous Solutions; Marcel Dekker: New York, 1985. Kissinger, P. T.; Heineman, W. R. Laboratory Techniques in Electroanalytical Chemistry; Marcel Dekker: New York, 1984. https://pineresearch.com/shop/products/potentiostats/wavenow-series/wavenow-wireless/. https://pineresearch.com/shop/kb/applications/general-electrochemistry/potentiostat-glovebox-installation/. Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Fundamentals and Applications, 2nd ed.; Wiley: New York, 2001. Treimer, S.; Tang, A.; Johnson, D. C. Electroanalysis 2002, 14 (3), 165–171. https://pineresearch.com/shop/kb/theory/hydrodynamic-electrochemistry/koutecky-levich-analysis/. Gunasingham, H. In Electrochemistry, Past and Present; Stock, J. T., Orna, M. V., Eds.; ACS Symposium Series American Chemical Society: Washington, DC, 1989; vol. 390. (Chapter 17). Zoski, C. G., Ed.; In Handbook of Electrochemistry; Elsevier: Amsterdam, The Netherlands, 2006. Graham, D. J. Standard Operating Procedures for Cyclic Voltammetry; https://sop4cv.com/index.html. Accessed June 2021. Flato, J. B. Anal. Chem. 1972, 44 (11), 75A–87A. Laborda, E.; González, J.; Molina, Á. Electrochem. Commun. 2014, 43, 25–30. Barker, G. C.; Jenkins, I. L. Analyst 1952, 77 (920), 685–696. Rudolph, M. J. Electroanal. Chem. 2003, 543 (1), 23–39. Feldberg, S. W. J. Am. Chem. Soc. 1966, 88 (3), 390–393. Britz, D.; Strutwolf, J. Digital Simulation in Electrochemistry; Springer: Berlin, New York, 2016. Bieniasz, L. K. Comput. Chem. 1992, 16 (1), 11–14. Bieniasz, L. K. mput. Chem. 1993, 17 (4), 355–368. Bieniasz, L. K. mput. Chem. 1997, 21 (1), 1–12. https://home.cyf-kr.edu.pl/nbbienia/elsim3ad.html Rudolph, M.; Reddy, D. P.; Feldberg, S. W. Anal. Chem. 1994, 66 (10), 589A–600A. http://www.digielch.de/. Gosser, D. K. Cyclic Voltammetry: Simulation and Analysis of Reaction Mechanisms; VCH, New York, NY: New York, NY, 1993. López, I.; Le Poul, N. Coord. Chem. Rev. 2021, 436, 213823–221844. Kuwana, T.; Darlington, R. K.; Leedy, D. W. Anal. Chem. 1964, 36 (10), 2023–2025. Bard, A. J. J. Chem. Educ. 1983, 60 (4), 302. Marcus, R. A. On the Theory of Oxidation-Reduction Reactions Involving Electron Transfer. I. J. Chem. Phys. 1956, 24 (5), 966–978. Johnson, M. D. In Comprehensive Organometallic Chemistry; Wilkinson, G., Stone, F. G. A., Abel, E. W., Eds.; Pergammon Press: Oxford, 1982; vol. 4; p 479. Gagne, R. R.; Koval, C. A.; Lisensky, G. C. Ferrocene as an Internal Standard for Electrochemical Measurements. Inorg. Chem. 1980, 19 (9), 2854–2855. Malischewski, M.; Adelhardt, M.; Sutter, J.; Meyer, K.; Seppelt, K. Science 2016, 353 (6300), 678–682.

282 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. 105. 106. 107. 108. 109. 110. 111. 112. 113. 114. 115. 116. 117. 118. 119. 120. 121. 122. 123. 124. 125.

Electrochemistry in Organometallic Chemistry Goodwin, C. A. P.; Giansiracusa, M. J.; Greer, S. M.; Nicholas, H. M.; Evans, P.; Vonci, M.; Hill, S.; Chilton, N. F.; Mills, D. P. Nat. Chem. 2021, 13 (3), 243–248. Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Fundamentals and Applications, 2nd ed.; Wiley: New York, 2001; pp 49–51. Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Fundamentals and Applications, 2nd ed.; Wiley: New York, 2001; pp 52–53. Hopkins Leseberg, J. A.; Lionetti, D.; Day, V. W.; Blakemore, J. D. Organometallics 2021, 40 (2), 266–277. Lionetti, D.; Day, V. W.; Lassalle-Kaiser, B.; Blakemore, J. D. Chem. Commun. 2018, 54 (14), 1694–1697. Bourrez, M.; Molton, F.; Chardon-Noblat, S.; Deronzier, A. Angew. Chem. Int. Ed. 2011, 50 (42), 9903–9906. Hartl, F.; Rossenaar, B. D.; Stor, G. J.; Stufkens, D. J. Recl. Trav. Chim. Pays-Bas 1995, 114 (11− 12), 565–570. Hopkins, J. A.; Lionetti, D.; Day, V. W.; Blakemore, J. D. Organometallics 2019, 38 (6), 1300–1310. See Supporting Information Figure S48 of publication for electrochemical control experiment to confirm the absence of a chloride in TBAPF6/CH3CN electrolyte. Lionetti, D.; Day, V. W.; Blakemore, J. D. Organometallics 2017, 36 (10), 1897–1905. Henke, W. C.; Lionetti, D.; Moore, W. N. G.; Hopkins, J. A.; Day, V. W.; Blakemore, J. D. ChemSusChem 2017, 10 (22), 4589–4598. Hopkins, J. A.; Lionetti, D.; Day, V. W.; Blakemore, J. D. J. Organomet. Chem. 2020, 921, 121294. Boyd, E. A.; Lionetti, D.; Henke, W. C.; Day, V. W.; Blakemore, J. D. Inorg. Chem. 2019, 58 (6), 3606–3615. Hershberger, J. W.; Klingler, R. J.; Kochi, J. K. J. Am. Chem. Soc. 1983, 105, 61. Grass, V.; Lexa, D.; Momenteau, M.; Savéant, J.-M. J. Am. Chem. Soc. 1997, 119 (15), 3536–3542. Goodridge, F.; King, C. In Techniques of Electroorganic Synthesis; Weinberg, N. L., Ed.; John Wiley: New York, 1974. Harrar, J. E. In Electroanalytical Chemistry; Bard, A. J., Ed.; Marcel Dekker: New York, 1975; vol. 8. Bard, A. J.; Zoski, C. G. Electroanalytical Chemistry: A Series of Advances: Volume 23; CRC Press, 2010. Shono, T. Electroorganic Synthesis; Academic Press, 1991. Kyriacou, D. K. Basics of Electroorganic Synthesis; Wiley-Interscience: New York, 1981. Fry, A. J. Synthetic Organic Electrochemistry; John Wiley & Sons, 1989. Kissinger, P. T.; Heineman, W. R. Laboratory Techniques in Electroanalytical Chemistry; Marcel Dekker: New York, 1984; pp 504–506. Robinson, J. R.; Gordon, Z.; Booth, C. H.; Carroll, P. J.; Walsh, P. J.; Schelter, E. J. J. Am. Chem. Soc. 2013, 135 (50), 19016–19024. Barr, J. L.; Kumar, A.; Lionetti, D.; Cruz, C. A.; Blakemore, J. D. Organometallics 2019, 38 (9), 2150–2155. Smieja, J. M.; Sampson, M. D.; Grice, K. A.; Benson, E. E.; Froehlich, J. D.; Kubiak, C. P. Inorg. Chem. 2013, 52 (5), 2484–2491. Chapovetsky, A.; Patel, P.; Liu, C.; Sattelberger, A. P.; Kaphan, D. M.; Delferro, M. Organometallics 2020, 39 (24), 4430–4436. Lewis, N. S.; Nocera, D. G. PNAS 2006, 103 (43), 15729–15735. International Energy Agency, Data and Statistics for World Electricity Generation by Source; . International Energy Agency, World Energy Outlook; . Roberts, J. A. S.; Bullock, R. M. Inorg. Chem. 2013, 52, 3823–3835. Kütt, A.; Leito, I.; Kaljurand, I.; Sooväli, L.; Vlasov, V. M.; Yagupolskii, L. M.; Koppel, I. A. J. Organomet. Chem. 2006, 71 (7), 2829–2838. Kaljurand, I.; Kütt, A.; Sooväli, L.; Rodima, T.; Mäemets, V.; Leito, I.; Koppel, I. A. Extension of the Self-Consistent Spectrophotometric Basicity Scale in Acetonitrile to a Full Span of 28 pKa Units: Unification of Different Basicity Scales. J. Org. Chem. 2005, 70 (3), 1019–1028. Schwesinger, R. Nachr. Chem. Tech. Lab. 1990, 38 (10), 1214–1226. Costentin, C.; Robert, M.; Savéant, J.-M.; Tatin, A. PNAS 2015, 112 (22), 6882–6886. Espenson, J. H. In Chemical Kinetics and Reaction Mechanisms; Espenson, J. H., Ed.; McGraw-Hill: New York, 1995. Matheu, R.; Neudeck, S.; Meyer, F.; Sala, X.; Llobet, A. ChemSusChem 2016, 9 (23), 3361–3369. Zott, M. D.; Garrido-Barros, P.; Peters, J. C. ACS Catal. 2019, 9 (11), 10101–10108. Ahmad, E.; Rai, S.; Padhi, S. K. Int. J. Hydrog. Energy 2019, 44 (31), 16467–16477. Clark, M. L.; Cheung, P. L.; Lessio, M.; Carter, E. A.; Kubiak, C. P. ACS Catal. 2018, 8 (3), 2021–2029. Liu, X.; Li, F.-F.; Peng, P.; Licht, G.; Licht, S. Eur. J. Inorg. Chem. 2020, 2020 (15–16), 1428–1436. Crabtree, R. H. Chem. Rev. 2015, 115 (1), 127–150. Whitesides, G. M.; Hackett, M.; Brainard, R. L.; Lavalleye, J. P. P. M.; Sowinski, A. F.; Izumi, A. N.; Moore, S. S.; Brown, D. W.; Staudt, E. M. Organometallics 1985, 4 (10), 1819–1830. Sconyers, D. J.; Blakemore, J. D. Distinguishing Between Homogeneous and Heterogeneous Hydrogen-Evolution Catalysis with Molecular Cobalt Complexes. Chem. Commun. 2017, 53 (53), 7286–7289. Lin, Y.; Finke, R. G. Inorg. Chem. 1994, 33 (22), 4891–4910. Sconyers, D. J.; Blakemore, J. D. Electrodeposition Behavior of Homoleptic Transition Metal Acetonitrile Complexes Interrogated With Piezoelectric Gravimetry. Analyst 2020, 145 (2), 466–477. Sconyers, D. J.; Blakemore, J. D. Distinguishing Deposition, Corrosion, and Stripping of Transient Heterogeneous Materials During Molecular Electrocatalysis. Dalton Trans. 2019, 48 (19), 6372–6382. Nie, W.; Wang, Y.; Zheng, T.; Ibrahim, A.; Xu, Z.; McCrory, C. C. L. ACS Catal. 2020, 10 (9), 4942–4959. REDOX n.d. This bulk electrolysis cell may be available for purchase at redox.me; item: bulk electrolysis basic cell – 50 mL. https://redox.me/products/bulk-electrolysis-basiccell-50-ml. Accesses June 2021. Harris, D. C. Quantitative Chemical Analysis; W.H. Freeman and Co: New York, NY, 2007. Print. Eaton, G. R.; Eaton, S. S.; Barr, D. P.; Weber, R. T. Quantitative EPR; Springer-Verlag: Vienna, 2010. Dau, H.; Limberg, C.; Reier, T.; Risch, M.; Roggan, S.; Strasser, P. ChemCatChem 2010, 2 (7), 724–761. Blakemore, J. D.; Crabtree, R. H.; Brudvig, G. W. Chem. Rev. 2015, 115 (23), 12974–13005. Hunter, B. M.; Gray, H. B.; Müller, A. M. Chem. Rev. 2016, 116 (22), 14120–14136. Yagi, M.; Kaneko, M. Chem. Rev. 2001, 101 (1), 21–36. Meyer, T. J.; Huynh, M. H. V.; Thorp, H. H. Angew. Chem. Int. Ed. 2007, 46 (28), 5284–5304. Armstrong, D. A.; Huie, R. E.; Koppenol, W. H.; Lymar, S. V.; Merényi, G.; Neta, P.; Ruscic, B.; Stanbury, D. M.; Steenken, S.; Wardman, P. Pure Appl. Chem. 2015, 87 (11–12), 1139–1150. Harriman, A.; Richoux, M.-C.; Christensen, P. A.; Mosseri, S.; Neta, P. J. Chem. Soc. Faraday Trans. 1987, 83 (9), 3001–3014. Harriman, A.; Thomas, J. M.; Millward, G. R. New J. Chem. 1987, 11, 757–762. Najafpour, M. M.; Ehrenberg, T.; Wiechen, M.; Kurz, P. Angew. Chem. Int. Ed. 2010, 49, 2233–2237. Gersten, S. W.; Samuels, G. J.; Meyer, T. J. J. Am. Chem. Soc. 1982, 104 (14), 4029–4030. Duan, L.; Bozoglian, F.; Mandal, S.; Stewart, B.; Privalov, T.; Llobet, A.; Sun, L. Nat. Chem. 2012, 4 (5), 418–423. Schulze, M.; Kunz, V.; Frischmann, P. D.; Würthner, F. Nat. Chem. 2016, 8 (6), 576–583. McDaniel, N. D.; Coughlin, F. J.; Tinker, L. L.; Bernhard, S. Cyclometalated Iridium(III) Aquo Complexes: Efficient and Tunable Catalysts for the Homogeneous Oxidation of Water. J. Am. Chem. Soc. 2008, 130, 210–217. Hull, J. F.; Balcells, D.; Blakemore, J. D.; Incarvito, C. D.; Eisenstein, O.; Brudvig, G. W.; Crabtree, R. H. J. Am. Chem. Soc. 2009, 131 (25), 8730–8731.

Electrochemistry in Organometallic Chemistry

283

126. Zhao, Y.; Yang, K. R.; Wang, Z.; Yan, X.; Cao, S.; Ye, Y.; Dong, Q.; Zhang, X.; Thorne, J. E.; Jin, L.; Materna, K. L.; Trimpalis, A.; Bai, H.; Fakra, S. C.; Zhong, X.; Wang, P.; Pan, X.; Guo, J.; Flytzani-Stephanopoulos, M.; Brudvig, G. W.; Batista, V. S.; Wang, D. PNAS 2018, 115 (12), 2902–2907. 127. Najafpour, M. M.; Renger, G.; Hołynska, M.; Moghaddam, A. N.; Aro, E.-M.; Carpentier, R.; Nishihara, H.; Eaton-Rye, J. J.; Shen, J.-R.; Allakhverdiev, S. I. Chem. Rev. 2016, 116 (5), 2886–2936. 128. Singh, A.; Spiccia, L. Coord. Chem. Rev. 2013, 257 (17), 2607–2622. 129. Blakemore, J. D.; Schley, N. D.; Balcells, D.; Hull, J. F.; Olack, G. W.; Incarvito, C. D.; Eisenstein, O.; Brudvig, G. W.; Crabtree, R. H. J. Am. Chem. Soc. 2010, 132 (45), 16017–16029. 130. Savini, A.; Bellachioma, G.; Ciancaleoni, G.; Zuccaccia, C.; Zuccaccia, D.; Macchioni, A. Chem. Commun. 2010, 46 (48), 9218–9219. 131. Blakemore, J. D.; Schley, N. D.; Olack, G. W.; Incarvito, C. D.; Brudvig, G. W.; Crabtree, R. H. Chem. Sci. 2011, 2 (1), 94–98. 132. Schley, N. D.; Blakemore, J. D.; Subbaiyan, N. K.; Incarvito, C. D.; D’Souza, F.; Crabtree, R. H.; Brudvig, G. W. J. Am. Chem. Soc. 2011, 133 (27), 10473–10481. 133. Adli, N. M.; Zhang, H.; Mukherjee, S.; Wu, G. J. Electrochem. Soc. 2018, 165 (15), J3130–J3147. 134. Dunn, P. L.; Cook, B. J.; Johnson, S. I.; Appel, A. M.; Bullock, R. M. J. Am. Chem. Soc. 2020, 142 (42), 17845–17858. 135. Habibzadeh, F.; Miller, S. L.; Hamann, T. W.; Smith, M. R. PNAS 2019, 116 (8), 2849–2853. 136. Zott, M. D.; Peters, J. C. J. Am. Chem. Soc. 2021, 143 (20), 7612–7616. 137. McKone, J. R.; Marinescu, S. C.; Brunschwig, B. S.; Winkler, J. R.; Gray, H. B. Chem. Sci. 2014, 5 (3), 865–878. 138. Tong, L.; Duan, L.; Zhou, A.; Thummel, R. P. Coord. Chem. Rev. 2020, 402, 213079. 139. Cheng, N.; Stambula, S.; Wang, D.; Banis, M. N.; Liu, J.; Riese, A.; Xiao, B.; Li, R.; Sham, T.-K.; Liu, L.-M.; Botton, G. A.; Sun, X. Nat. Commun. 2016, 7 (1), 13638. 140. Jaramillo, T. F.; Jørgensen, K. P.; Bonde, J.; Nielsen, J. H.; Horch, S.; Chorkendorff, I. Science 2007, 317 (5834), 100–102. 141. Valdez, C. N.; Dempsey, J. L.; Brunschwig, B. S.; Winkler, J. R.; Gray, H. B. PNAS 2012, 109 (39), 15589–15593. 142. Michael Gratzel, U. K. Chem. Ber. 1989, 122, 1869–1880. 143. Francke, R.; Schille, B.; Roemelt, M. Chem. Rev. 2018, 118 (9), 4631–4701. 144. Qiao, J.; Liu, Y.; Hong, F.; Zhang, J. Chem. Soc. Rev. 2014, 43 (2), 631–675. 145. Frese, K. W. Electrochemical Reduction of Carbon Dioxide at Solid Electrodes. In Electrochemical and Electrocatalytic Reactions of Carbon Dioxide; Sullivan, B. P., Krist, K., Guard, H. E., Eds.; Elsevier: Amsterdam, 2012;; pp 145–216. 146. Bhugun, I.; Lexa, D.; Savéant, J.-M. J. Am. Chem. Soc. 1996, 118 (7), 1769–1776. 147. Hawecker, J.; Lehn, J.-M.; Ziessel, R. J. Chem. Soc. Chem. Commun. 1984, (6), 328–330. 148. Sullivan, B. P.; Bolinger, C. M.; Conrad, D.; Vining, W. J.; Meyer, T. J. J. Chem. Soc. Chem. Commun. 1985, 20, 1414–1416. 149. Tignor, S. E.; Kuo, H.-Y.; Lee, T. S.; Scholes, G. D.; Bocarsly, A. B. Organometallics 2019, 38 (6), 1292–1299. 150. Henke, W. C.; Otolski, C. J.; Moore, W. N. G.; Elles, C. G.; Blakemore, J. D. Inorg. Chem. 2020, 59 (4), 2178–2187. 151. Machan, C. W.; Sampson, M. D.; Chabolla, S. A.; Dang, T.; Kubiak, C. P. Organometallics 2014, 33 (18), 4550–4559. 152. Sampson, M. D.; Nguyen, A. D.; Grice, K. A.; Moore, C. E.; Rheingold, A. L.; Kubiak, C. P. J. Am. Chem. Soc. 2014, 136 (14), 5460–5471. 153. Bazhenova, T. A.; Shilov, A. E. Coord. Chem. Rev. 1995, 144, 69–145. 154. Nishibayashi, Y. Inorg. Chem. 2015, 54, 9234–9247. 155. Fields, S. Environ. Health Perspect. 2004, 112, A556–A563. 156. Chalkley, M. J.; Drover, M. W.; Peters, J. C. Chem. Rev. 2020, 120 (12), 5582–5636. 157. Del Castillo, T. J.; Thompson, N. B.; Peters, J. C. J. Am. Chem. Soc. 2016, 138 (16), 5341–5350. 158. Chalkley, M. J.; Garrido-Barros, P.; Peters, J. C. Science 2020, 369 (6505), 850–854. 159. Chalkley, M. J.; Del Castillo, T. J.; Matson, B. D.; Roddy, J. P.; Peters, J. C. ACS Cent. Sci. 2017, 3 (3), 217–223. 160. Horn, E. J.; Rosen, B. R.; Baran, P. S. ACS Cent. Sci. 2016, 2 (5), 302–308. 161. Francke, R.; Little, R. D. Chem. Soc. Rev. 2014, 43 (8), 2492–2521. 162. Siu, J. C.; Fu, N.; Lin, S. Acc. Chem. Res. 2020, 53 (3), 547–560. 163. Faraday, M. Ann. Phys. 1834, 109 (23–30), 433–451. 164. Kolbe, H. J. Prakt. Chem. 1847, 41 (1), 137–139. 165. Schöenbein, C. F. Liebigs Ann. Chem. 1845, 54, 164. 166. Yan, M.; Kawamata, Y.; Baran, P. S. Chem. Rev. 2017, 117 (21), 13230–13319. 167. Caix, C.; Chardon-Noblat, S.; Deronzier, A.; Moutet, J.-C.; Tingry, S. J. Organomet. Chem. 1997, 540 (1), 105–111. 168. Hou, Z.-W.; Mao, Z.-Y.; Zhao, H.-B.; Melcamu, Y. Y.; Lu, X.; Song, J.; Xu, H.-C. Angew. Chem. Int. Ed. 2016, 55 (32), 9168–9172. 169. Choi, G. J.; Zhu, Q.; Miller, D. C.; Gu, C. J.; Knowles, R. R. Nature 2016, 539 (7628), 268–271. 170. Tafel, J. Z. Phys. Chem. 1905, 50A, 641–712. 171. Li, D.; Lin, C.; Batchelor-McAuley, C.; Chen, L.; Compton, R. G. J. Electroanal. Chem. 2018, 826, 117–124. 172. Costentin, C.; Drouet, S.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 2012, 134, 11235–11242. 173. Sconyers, D. J.; Shaughnessy, C. I.; Lee, H.-J.; Subramaniam, B.; Leonard, K. C.; Blakemore, J. D. ChemSusChem 2020, 13, 6338–6345. 174. Pegis, M. L.; Roberts, J. A. S.; Wasylenko, D. J.; Mader, E. A.; Appel, A. M.; Mayer, J. M. Inorg. Chem. 2015, 54 (24), 11883–11888.

1.10

Organometallic Photosensitizers

Thomas S Teets and Yanyu Wu, Department of Chemistry, University of Houston, Houston, TX, United States © 2022 Elsevier Ltd. All rights reserved.

1.10.1 Introduction 1.10.1.1 Excited state processes 1.10.1.2 Electron transfer photosensitization 1.10.1.2.1 Photoredox catalysis 1.10.1.3 Energy transfer photosensitization 1.10.1.3.1 Energy transfer initiated photocatalysis 1.10.2 Zirconium photosensitizers 1.10.3 Group 6 photosensitizers 1.10.4 Group 8 photosensitizers 1.10.5 Iridium photosensitizers 1.10.5.1 Dimeric cyclometalated iridium complexes 1.10.5.2 Homoleptic tris-cyclometalated iridium complexes 1.10.5.3 Cationic bis-cyclometalated iridium complexes 1.10.5.3.1 Complexes with bipyridine-derived ancillary ligands 1.10.5.3.2 Carboxy-substituted complexes for solar cells 1.10.5.3.3 Cationic complexes with quinoline-derived cyclometalating ligands 1.10.5.3.4 Cationic bis-cyclometalated iridium photocatalytic dyads 1.10.5.3.5 Complexes with alternative diimine ancillary ligands: enhanced light absorption and charge separation 1.10.5.3.6 Alternative cyclometalating ligands for enhanced light absorption and charge-transfer lifetimes 1.10.5.4 Charge-neutral heteroleptic bis-cyclometalated iridium complexes 1.10.5.4.1 Complexes with acac ancillary ligands 1.10.5.4.2 Bis-cyclometalated iridium complexes with 2-picolinate ancillary ligands 1.10.5.4.3 Bis-cyclometalated iridium complexes with electron-rich ancillary ligands 1.10.5.5 Bis-cyclometalated iridium complexes with N-heterocyclic carbene (NHC)-derived ancillary ligands 1.10.5.6 Bis-cyclometalated iridium complexes for enantioselective photoredox transformations 1.10.5.7 Bis-cyclometalated iridium complexes with monodentate ancillary ligands 1.10.5.8 Cyclometalated iridium complexes with tridentate ligands 1.10.5.9 Summary and outlook 1.10.6 Rhodium photosensitizers 1.10.7 Palladium photosensitizers 1.10.8 Platinum photosensitizers 1.10.9 Coinage metal photosensitizers 1.10.10 Summary and conclusions Acknowledgments References

285 285 286 287 287 288 289 290 291 293 293 293 295 295 297 298 300 301 305 308 308 313 317 319 322 325 327 329 330 331 332 333 334 334 334

Nomenclature acac bpy C^N DSSC EnT Fc LC LLCT LMCT MLCT piq ppy pq PS SET

284

Acetylacetonate 2,20 -Bipyridine Cyclometalating Dye-sensitized solar cell Energy transfer Ferrocene Ligand-centered Ligand-to-ligand charge transfer Ligand-to-metal charge transfer Metal-to-ligand charge transfer 1-Phenylisoquinoline 2-Phenylpyridine 2-Phenylquinoline Photosensitizer Single-electron transfer

Comprehensive Organometallic Chemistry IV

https://doi.org/10.1016/B978-0-12-820206-7.00008-1

Organometallic Photosensitizers

e l FPL FD

1.10.1

285

Molar absorptivity Wavelength Photoluminescence quantum yield Singlet oxygen quantum yield

Introduction

Photosensitizers are photoactive chemicals that can absorb the energy from light, typically in the UVA or visible light region, to reach their excited states which then facilitate energy transfer or electron transfer processes. They are fundamental components in a variety of important applications including generation of solar fuels,1 photovoltaics,2 photocatalysis,3,4 and chemical sensors,5,6 among others. Photosensitizers are typically either organic dyes,7–9 semiconductors,10 polymers,11 or metal complexes.12,13 Metal coordination complexes are especially prominent, due to the ability to control their photophysical and electrochemical attributes by changing the identity of the metal centers and the organic ligands. In this article, we introduce the general photoinduced excited state processes, important photophysical and electrochemical properties, as well as electron-transfer and energy-transfer pathways of metal-based molecular photosensitizers. We will also highlight some classical and recently emerged organometallic photosensitizers, mainly focusing on those that have been applied in photocatalysis and dye-sensitized photovoltaics. We limit the scope to transition metal photosensitizers that fit the traditional definition of “organometallic,” including at least one metal-carbon bond, usually in the form of one or more cyclometalated aryl rings from the supporting ligands. As a result, we will not describe other prominent classes of molecular photosensitizers, and refer the reader to other recent review articles and books describing the important ruthenium polypyridyl14–19 and organic9,20,21 photosensitizer classes. Furthermore, while there are many luminescent organometallic complexes that could potentially be described and used as photosensitizers, we limit our discussions primarily to compounds which have actually been used as sensitizers, in processes involving photoinduced electron and/or energy transfer. Thus, there is a large suite of phosphorescent organometallic complexes, which have been very successful in electroluminescent applications,22–24 which are not covered in this article unless they have also been investigated as photosensitizers. The majority of prominent organometallic photosensitizers are iridium-based and thus will be heavily featured in this article, but we aim to provide comprehensive coverage across the transition-metal series. In the first section of this article, we provide an overview of the concepts and terminology that are important for evaluating and characterizing photosensitizers.

1.10.1.1

Excited state processes

Molecular organometallic photosensitizers typically consist of a transition metal center bonded to organic ligands possessing low-lying p orbitals. In most cases the highest occupied molecular orbitals (HOMOs) of these molecules are dominated by metal-based dp orbitals while their lowest unoccupied molecular orbitals (LUMOs) are centered on the p orbital. The excited state processes of these complexes can be illustrated by the simplified Jablonski diagram shown in Fig. 1. Upon absorption of light with specific wavelengths in the UV-Vis region, the photosensitizer (PS) is excited and converted from its ground state (S0) to the singlet excited state (S1), forming the excited species (1 PS). Transitions between the metal-centered orbitals and the ligand-centered orbitals, including the HOMO !LUMO transition, are often referred to as metal-to-ligand charge transfer (MLCT), and the corresponding singlet excited state is denoted as 1MLCT. The visible absorption bands in organometallic photosensitizers normally involve population of 1MLCT states. This MLCT process also results in formal oxidation of the metal center and reduction of the ligand. Ligand centered (LC) transitions from p to p orbital of the organic ligand give rise to the 1LC excited states. These 1LC states are normally higher in energy and thus excited in the UV but depending on the nature of the organometallic complex the lowest-energy S1 excited state can consist of a mixed configuration involving the 1MLCT and 1LC states. In principle the energy of the photoexcited singlet molecule can be released through fluorescence emission or nonradiative internal conversion to relax back to the ground state, but in complexes with heavy transition metals the lowest energy triplet excited state (T1) can be populated through efficient intersystem crossing (ISC). The T1 state, which again normally has 3MLCT and/or 3LC character, will then relax in the form of phosphorescence or thermally relax through nonradiative deactivation. In the case of highly efficient triplet

Fig. 1 Jablonski diagram depicting the excited state process of molecular photosensitizers.

286

Organometallic Photosensitizers

photosensitizers, ISC is very fast (on the orders of fs or ps), which bypasses fluorescence and singlet-state nonradiative pathways. The triplet state species (3 PS) is the key component that participates in the following sensitization-initiated electron transfer and energy transfer, thus evaluating the photophysical and electrochemical properties of the triplet excited state is of particular importance. Having outlined the photoinduced excited state processes of molecular photosensitizers, it is instructive to discuss some of the related fundamental photophysical properties. Different applications have specific requirements for the photophysics of the photosensitizer. For applications such as dye-sensitized solar cells and solar fuels, it is desirable for the photosensitizer to absorb a broad range of solar radiation with high extinction coefficient (e). This is because their efficiency is proportional to the absorbed photons. For other applications such as organic photoredox catalysis and molecular sensing, the absorption range is less important. However, it is still favorable for the photosensitizer to be excited by visible light since the use of higher-energy UV photons can compromise the selectivity and may lead to photodegradation. Typical MLCT photosensitizers feature weak to moderate absorption (e < 15,000 M−1 cm−1) and broad absorption bands.25 The absorption intensity of these molecules can be improved by incorporating chromophores in the coordination center such as boron dipyrromethene (BODIPY), which demonstrates strong absorption in the visible light spectrum.26 The MLCT absorption energy can be shifted to longer wavelengths by minimizing the HOMO-LUMO energy gap, by destabilizing the metal dp orbital and/or lowering the ligand-based p orbital energy. Another important photophysical feature to consider when evaluating the photosensitizer is the excited state lifetime (t). Triplet photosensitizers normally have long-lived excited states as a result of the spin-forbidden electronic transition between the triplet excited state and the singlet ground state. Although a long-lived T1 state is detrimental to luminescence performance and optoelectronic applications, it is requisite for bimolecular electron-transfer and energy-transfer processes in photochemical transformations to outpace nonradiative deactivation. Designing molecular photosensitizers with long-lived excited states have been a research topic of great interest especially in the area of photocatalysis. The t can be prolonged by increasing the steric bulkiness of the molecule to restrict relaxation pathways that involve excited state distortion.27 Another common strategy to increase the t is by maximizing the energy gap between the 3MLCT state and ligand-field metal-centered triplet excited state (3MC). This is because thermal population of the 3MC state by the 3MLCT state when they are close in energy contributes to detrimental nonradiative deactivation. This has been achieved by increasing the p-conjugation of the ligand to lower the 3MLCT energy or introducing electron-donating ancillary ligand to destabilize the 3MC state.28 This same method can also engender higher phosphorescence quantum yields. It has also been shown that changing the nature of the luminescent excited state from 3MLCT to 3LC can lead to elongation of the triplet excited state lifetime.29

1.10.1.2

Electron transfer photosensitization

After reaching the excited state, in the presence of a suitable reaction partner the photosensitizer can engage in dynamic quenching through either electron transfer or energy transfer pathways. “Quenching” is defined as a physical conversion or change in molecular state that reduces the luminescence of the excited state molecule and it is a fundamental step in sensitization. During the electron-transfer process, also known as single-electron transfer (SET), the photosensitizer can function as an electron donor or acceptor depending on the relative redox potentials of the  PS and the quencher. The quencher can be a reaction substrate, catalyst or semiconductor depending on the application and the preferred mechanism of the targeted chemical reaction. When the photosensitizer donates an electron to the quencher, it is classified as a photoreductant, and the quenching event results in the formation of PS+, the one-electron oxidized form of the photosensitizer. Alternatively, when the photosensitizer receives an electron from the quencher it will be described as a photooxidant, which results in the formation of PS−, the one-electron reduced form of the photosensitizer. The excited state redox properties can be defined by the excited state redox potentials calculated from Eqs. 1 and 2. The excited state energy (E0,0) refers to the energy difference between the lowest vibrational levels of the electronic ground state and the excited state, which can also be abbreviated as ET1 for the lowest triplet excited state. This value can be estimated in a few different ways from UV-vis absorption and photoluminescence spectroscopy or determined computationally. Complexes in their excited states are generally more reducing and oxidizing than in their ground states, with a larger value of E0,0 corresponding to both stronger reducing and oxidizing ability. The drawback is that the E0,0 value represents the minimum amount of energy needed to reach the photoactive excited state, so a higher E0,0 also implies that a larger amount of energy, i.e. a shorter wavelength of light, is required to activate the photosensitizer, which can be a limitation in certain applications. The excited state redox potentials [E(PS+/ PS)] and [E( PS/PS− )] are determined both by the E0,0 and the standard ground-state redox potential. The ground-state potentials E(PS+/PS) and E(PS/PS−) are obtained from the half-wave potentials (E1/2) of the first one-electron oxidation and reduction of the photosensitizer, respectively. A more negative value of E(PS+/ PS) corresponds to higher reducing power (stronger photoreductant), while a more positive value of E( PS/PS−) demonstrates higher oxidizing power (stronger photooxidant). As a note about terminology commonly used in this article and in the broader literature, for metal-based photosensitizers the potentials are often abbreviated in terms of the formal oxidation states of the metal. Thus, for an Ir(III) photosensitizer E(PS+/ PS) would be commonly written as E(IrIV/ IrIII), and E( PS/PS−) would be commonly written as E( IrIII/IrII). In either case, these potentials refer to the one-electron redox processes that occur from the photosensitizer’s excited state. Photoreductant : EðPS + = PSÞ ¼ EðPS + =PSÞ 

E0, 0

Photooxidant : Eð PS=PS Þ ¼ E0, 0 + EðPS=PS Þ −

−

(1) (2)

Organometallic Photosensitizers

1.10.1.2.1

287

Photoredox catalysis

Most photocatalytic transformations proceed via the SET mechanism and these catalytic reactions are also known as photoredox catalysis. Such transformations are important in organic synthesis, polymer synthesis, photocatalyzed hydrogen production from water splitting, and photochemical CO2 reduction, among others.3,30,31 Fig. 2 displays the two major mechanistic cycles of SET photocatalysis, which are known as oxidative quenching and reductive quenching. In the oxidative quenching process, the excited state photosensitizer ( PS) is used as a photoreductant and donates an electron to an electron acceptor (EA), resulting in the generation of the radical anion of the EA (EA•−) and the one-electron oxidized species of the photosensitizer (PS+). The anionic EA•− is generally more reactive than its neutral congener, can initiate the catalytic reaction or in some cases a radical chain process, and normally cannot be generated through direct UV-Vis light absorption. The photosensitizer can be regenerated in the presence of an electron donor (ED). The EA and ED in this mechanism will be described as the oxidative quencher and sacrificial reductant, respectively. In the reductive quenching pathway, the photosensitizer acts as a photooxidant, which receives an electron from an electron donor (ED), forming the one-electron reduced species (PS−) and the cationic radical of the ED (ED•+). The reductive cycle then closes as the PS– is oxidized to its original ground state in the presence of a sacrificial oxidant reagent. These two SET pathways signify the importance of both the excited state and electronic ground state redox properties of the photosensitizer. The excited-state SET process is governed by the excited state redox potential while the ground-state redox potentials are significant for the subsequent reactions of PS+ or PS− and eventual regeneration of the photosensitizer. Because of these thermodynamic considerations, the choice of suitable oxidative/reductive quenchers and sacrificial redox reagents are also essential to optimize a photoredox transformation. In practice it is usually not possible to predict the electron-transfer mechanism of a photoredox reaction a priori, and there are a myriad of other factors beyond redox potentials that determine the best choice of sacrificial reagents. In addition, the reversibility of the ground state redox events is necessary to ensure high turnover numbers of the photosensitizer and avoid decomposition pathways of PS+ or PS−.

Fig. 2 General mechanistic pathways of oxidative and reductive quenching processes.

1.10.1.3

Energy transfer photosensitization

Photosensitizers can also transfer their excited energy to an energy acceptor (EnA), resulting in the formation of the excited state of the latter. There are two main proposed mechanisms for light-induced energy transfer (EnT), namely the Förster resonance energy transfer and the Dexter energy transfer as shown in Fig. 3. Both EnT mechanisms result in electronic transition of the  PS back to its ground state while the EnA is promoted to its excited state ( EnA). Förster energy transfer is induced by nonradiative Coulombic (dipole-dipole) resonance interaction between the 1 PS and EnA, i.e. the 1 PS induces a dipole in the ground state of the EnA by charge repulsion.32–34 Such interaction occurs when there is significant overlap between the emission spectrum of the photosensitizer and the absorption spectrum of the acceptor. The Förster transfer strongly depends on the distance between the two molecules and such mechanism has found applications such as bioimaging, fluorescence labelling and protein interaction.35–37 Förster resonance energy transfer does not apply to triplet sensitizers since it requires the electronic transition of a triplet excited state to the singlet ground state via the dipole-dipole mechanism, which is forbidden by Wigner’s spin conservation law.38 Thus, they most commonly employ the Dexter energy transfer mechanism. The Dexter mechanism involves electron exchange interaction between the electron in the excited state of the photosensitizer and that in the ground state of the acceptor.39 It can be described as an electron in the LUMO of the  PS being donated to the unoccupied LUMO of the EnA, while simultaneously the HOMO of the  PS receives an electron from the EnA. This electronic exchange interaction requires significant molecular orbital overlap between these two molecules and occurs when they are in close proximity (1.5 ms) and quantum yields (>0.40) that are over an order of magnitude larger than W2, when measured in toluene. On account of the lower excited-state energies these compounds are not as potent photoreductants as W2 (E(WI/ W0) −2.8 V vs. Fc+/Fc for the oligoaryl isocyanide complexes), but they are still capable of generating ketyl radicals via photoinduced electron transfer to benzophenone and acetophenone. Wenger et al. have introduced related group 6 complexes with chelated bis(isocyanides), which are expected to be more robust than their monodentate analogues. These Mo(0) and Cr(0) complexes (Mo3, Mo4, and Cr3) with chelating bidentate isocyanide ligands are likewise strong photoreductants upon visible light excitation.48–50 Mo3 has an excited-state redox potential E(MoI/ Mo0) ¼ −2.6 V vs Fc+/Fc, not as strong as the tungsten isocyanide complexes described above but still capable of challenging reduction reactions. The Mo complex Mo3 can facilitate the rearrangement of acyl-cyclopropane to 2,3-dihydrofuran, which typically occurs under a strong reducing environment.49 The chromium complex Cr3 has not been used as a photosensitizer for catalytic reactions, but it is a rare example of a first-row transition metal complex with room-temperature luminescence and >ns lifetime (2.2 ns), and it was shown to be capable of triplet energy transfer to anthracene.50 In their most recent work which compared Mo3 and Mo4, the same group showed that the tert-butyl-substituted analogue Mo4 has nearly identical UV-vis absorption and emission wavelengths, albeit with higher molar absorptivity in the visible region and a lifetime that is at least an order of magnitude longer.48 They showed in this work that these Mo photosensitizers are capable of photoredox activation of aryl iodide and aryl bromide substrates, with complex Mo4 having the best photoreactivity and promoting intramolecular radical cyclization reactions of several aryl iodide substrates, where the cyclization occurs via radical addition of the photogenerated aryl radical to a second tethered aryl or heteroaryl radical. These group 6 photosensitizers offer the promise of very reducing excited-state potentials, but at this point their applications in photoredox catalysis are nascent. It is still an open question whether they will prove to be as robust and versatile as other d6 photosensitizers, e.g. Ru(II) and Ir(III). All of the examples presented in Fig. 7 are homoleptic compounds with isocyanide ligands, which is partly out of necessity since p-accepting ligands like isocyanides are best suited to stabilize the low M(0) oxidation

Organometallic Photosensitizers

291

Fig. 7 Group 6 isocyanide photosensitizers.

states in these complexes. Further tuning of the excited states of these compounds is in principle possible with more complex ligand combinations, but it is not obvious whether the synthetic chemistry of these compounds can be adapted for such modification. That said, there is no particular reason why photosensitizers in other oxidation states with other ligand classes couldn’t be developed, in particular for tungsten where higher oxidation state W(II) and W(IV) complexes with organometallic functionalities are common, but with little known about their photochemistry. Given the early success with the tungsten photosensitizers described here, there is certainly an opportunity for exploration of other group 6 structures types which may likewise have long-lived excited states capable of charge transfer.

1.10.4

Group 8 photosensitizers

Iron complexes are considered as desirable candidates for photoredox catalysis and other photosensitizer applications due to the large abundance and rich coordination chemistry of iron. One of the main challenges when designing iron-based photosensitizers is their low-lying metal-centered (ligand-field) states that deactivate the desired charge-transfer states, leading to short-lived excited states. Studies concerning Fe-based photosensitizers have been focusing on stabilizing the charge-transfer states and/or destabilizing the ligand-field states. A significant breakthrough is marked by the complex Fe(phtmeimb)+2 [Fe1, phtmeimb ¼ phenyl(tris(3methylimidazol-1-ylidene)borate)], which features a strong s-donor tris-carbene ligand designed to destabilize the ligand-field states (Fig. 8). This complex exhibits photoluminescence from a 2LMCT state with a lifetime of 2.0 ns and a quantum yield of 2% at room temperature.51 This initial study showed that this compound is capable of bimolecular photoinduced electron transfer

292

Organometallic Photosensitizers

Fig. 8 Group 8 organometallic photosensitizers.

reactions, reacting with diphenylamine via reductive quenching and methyl viologen via oxidative quenching. A much more detailed study of the excited-state charge-transfer kinetics of Fe1 appeared more recently, investigating reductive quenching processes involving triethylamine and N,N-dimethylaniline electron donors.52 The second-row group 8 transition metal, ruthenium, is ubiquitous in metal-based photosensitizers, with over 50 years of research on the [Ru(bpy)3]2+ (bpy ¼ 2,20 -bipyridine) family of complexes.53 These ruthenium polypyridyl complexes have been especially prominent in dye-sensitized solar cells14 and more recently in photoredox catalysis.3 The vast majority of compounds from this family are supported by nitrogen-donor chelating ligands, but there have been some developments of organometallic ruthenium photosensitizers that have one or more ruthenium-carbon bonds. Recently Ru polypyridyl complexes containing cyclometalating ligands have received increasing attention in regard to their photophysical characteristics. The chemical structures of several organometallic Ru photosensitizers are outlined in Fig. 8. Pioneering work reported by Grätzel et al. shows a novel molecular Ru chromophore with a 2-(2,4-difluorophenyl)pyridine ligand, (Ru1), which surpasses the benchmark power conversion efficiency of 10% when applied as a photosensitizer in dye-sensitized solar cell (DSSC) devices.54 The presence of fluorine substituents is the key to its high efficiency by tuning the basicity of the cyclometalating ligand to adjust the ground state redox potential. Its remarkable performance is also attributed to the complex’s broad range of visible light absorption and high molar extinction coefficient due to the strong electron-donating cyclometalating ligand. The absorption spectra of polypyridyl complexes of Ru are characterized by strong intraligand (IL) p !p transitions in the UV region and MLCT transitions in the visible region. Isoelectronic replacement of the neutral nitrogen atom in one of the polypyridine ligands by a s-donating carbanionic atom generally engenders stronger electron donation to the Ru center and induces significant changes in their photophysics. Koten et al. and Berlinguette et al. reported a series of cyclometalated Ru complexes and elucidated

Organometallic Photosensitizers

293

their fundamental photophysical and electrochemical properties.55–59 Fig. 8 shows some of these Ru complexes (Ru2–Ru6), which bear C^N, C^N^N or N^C^N cyclometalating ligands.57,60 Generally, these complexes exhibit MLCT transitions with bathochromic shifts with respect to their analogues supported by polypyridine or phenanthroline ligands with only N-donors. This is because the cyclometalating ligands destabilize the Ru-centered HOMOs while their LUMOs are still exclusively dominated by the polypyridine motifs. The introduction of an organometallic RudC bond also reduces the symmetry of the molecules resulting in a broader range of absorption in the visible region. These characteristics are especially important in solar-to-energy conversion applications since conventional homoleptic Ru polypyridyl complexes generally absorb relatively little solar energy at longer wavelengths. Ru complexes containing neutral N-heterocyclic carbenes (NHCs) have been investigated recently for their photophysical properties and a few examples of these complexes were studied in photosensitizing applications. Complex Ru7 featuring a NHC-pyridine ligand achieves MLCT absorption up to 750 nm attributable to the strong NHC s-donor and demonstrates excellent efficiency in a DSSC,  10%.61 The homoleptic Ru complex with tridentate NHC-pyridine ligands, Ru8, shows an excited-state potential [E(Ru3+/ Ru2+)] of −1.61 V vs. Fc+/Fc that corresponds to strong reducing power, and an exceptionally long excited-state lifetime of 0.82 ms in acetonitrile and 3.1 ms in water (Br− was used as the counteranion) that are significantly longer than the 62 0 0 00 This complex can sensitize the photoreduction of water for hydrogen gas analogous Ru(tpy)2+ 2 (tpy ¼ 2,2 :6 ,2 -terpyridine). production in the presence of [Co(bpy)3]Cl2 (bpy ¼ 2,20 -bipyridine) as the co-catalyst and triethanolamine as the sacrificial electron donor without the use of an electron delay.63

1.10.5

Iridium photosensitizers

1.10.5.1

Dimeric cyclometalated iridium complexes

Among the earliest cyclometalated iridium complexes to be extensively studied have the general formula [Ir(C^N)2(m-Cl)]2, where C^N is a cyclometalating ligand consisting of a nitrogen heterocycle and orthometalated aryl group, most commonly 2-phenylpyridine or a closely-related derivative. These compounds are easy to prepare, forming as major products when IrCl3
nH2O and the respective protonated cyclometalating ligand precursor are heated in high-boiling alcohol solvents like 2-ethoxyethanol (Scheme 1).64 Watts and co-workers were the first to systematically study the photophysical and redox properties of compounds of this type.65,66 They prepared three variants, [Ir(ppy)2(m-Cl)]2 (ppy ¼ 2-phenylpyridine) and two methyl-substituted variants (Scheme 1). They showed that the methyl substituents in the substituted variants have a modest effect on the formal IrIV/IrIII redox couple, which shifts cathodically by 100 mV relative to the unsubstituted ppy dimer, and induces a small red shift of the emission, 0.04 eV when the ppy phenyl ring is methylated (Ir2), and 0.08 eV when the pyridine ring has the methyl group (Ir3). The combined photophysical and redox properties of these compounds indicate that they are strong photoreductants, with excited-state redox potentials, E(IrIV/ IrIII), of ca. −2.1 V vs. Fc+/Fc in each case. However, neither this early study nor any subsequent studies have further elaborated the excited-state redox chemistry or photosensitizer applications of these dimeric cyclometalated iridium complexes.

Scheme 1 Synthesis of dimeric cyclometalated iridium complexes.

1.10.5.2

Homoleptic tris-cyclometalated iridium complexes

Tris-cyclometalated iridium(III) complexes of the general formula Ir(C^N)3, have become among the most widely studied and broadly applicable classes of organometallic photosensitizers. These complexes can exist in the fac or mer geometry, with the former being the thermodynamically stable product and the one that has been more widely studied. They were also originally introduced by Watts, and initially were isolated as minor side products (10% or less) during the synthesis of chloride-bridged cyclometalated iridium dimers using the conditions shown in Scheme 1.67 A more rational synthesis, described in Scheme 2, starts with the iridium(III) precursor Ir(acac)3 (acac ¼ acetylacetonate), which is treated with an excess amount of the cyclometalating ligand

294

Organometallic Photosensitizers

precursor and heated at reflux in glycerol at 290  C.68 This procedure produces the homoleptic fac-Ir(C^N)3 products in moderate to good isolated yields of 40–75%, following chromatographic purification. Studies on the parent complex fac-Ir(ppy)3 (Ir4) showed that it is a potent photoreductant, with E(IrIV/ IrIII) ¼ −2.1 V vs. Fc+/Fc, with the excited state efficiently quenched by electron-acceptor quenchers via rapid excited-state energy transfer.67 Quenching studies were not performed on the substituted variants shown in Scheme 2, nevertheless all of Ir5–Ir11 have similar excited-state lifetimes spanning 1.9–2.9 ms, and while the substituents have modest impacts on the ground-state IrIV/IrIII redox couple and triplet excited-state energies, all complexes have rather similar excited-state redox potentials and are expected to be potent photoreductants.

Scheme 2 Synthesis of homoleptic fac-Ir(C^N)3 complexes.

Homoleptic fac-Ir(C^N)3 complexes have become ubiquitous in the field of photoredox organic synthesis, particularly when using substrates that are challenging to reduce. In most cases the complex Ir4 is used, with many successful photoredox methodologies in small-molecule organic and polymer synthesis employing this unsubstituted “parent” photosensitizer.3,69–76 In this context there has also been some recent effort on preparing new substituted fac-Ir(ppy)3 variants for photocatalytic transformations. Though these modifications do not have large impacts on the excited-state redox potentials of the complexes, they are sometimes critical for reaction discovery. As part of their optimization of aldehyde a-benzylation reactions using alcohols, Nacsa and MacMillan prepared and screened nine fac-Ir(ppy)3 variants.77 Their screen included the known complexes Ir4, Ir7, Ir10, and Ir11, as well as the newly-synthesized derivatives Ir12–Ir16 summarized in Fig. 9. In this study they examined photocatalyst structure-activity relationships, measuring their performance in the photocatalytic a-benzylation of hydrocinnamaldehyde with 4-(hydroxymethyl)pyridine, focusing on the selectivity for the desired a-benzyl aldehyde product over the undesired

Fig. 9 Summary of some recently developed fac-Ir(ppy)3 derivatives.

Organometallic Photosensitizers

295

4-methylpyridine reduction product. They found a good correlation between product selectivity and the ground-state IrIV/IrIII redox potential of the photosensitizer, which led to the conclusion that the IrIV state that forms from oxidative quenching is involved in a deleterious oxidation step that results in the formation of the undesired side product. This work reveals that it is sometimes ground-state potentials, and not excited-state potentials, that are a key determinant of reaction outcomes in photoredox catalysis, necessitating careful consideration of all redox potentials, among many other possible factors, when optimizing a transformation. Weaver’s group has introduced an alternative to the synthetic procedure shown in Scheme 2, replacing the Ir(acac)3 precursor with the cheaper alternative IrCl3
nH2O,78 and using this modified approach to prepare Ir4, Ir7, Ir8, Ir9, and Ir14, as well as two other derivatives, Ir17 and Ir18, summarized in Fig. 9. Their work has shown that the substituent pattern on the fac-Ir (ppy)3 derivative can influence the E/Z selectivity of photocatalytic alkenylation reactions, with “small” catalysts like Ir4 and Ir8 giving Z-selectivity, and “large” catalysts like Ir7 giving primarily E stereochemistry for the product.79 This trend occurs because E/Z isomerization proceeds by a triplet energy transfer mechanism, which has a strong distance dependence and only occurs when close approach of the photosensitizer and catalyst is possible. Finally, Wenger’s group has prepared the water-soluble sulfonated facIr(ppy)3 derivatives Ir19–Ir22. Two-photon excitation of Ir19 produces hydrated electrons capable of performing a wide range of extreme reductions,80,81 and this series has also been used for visible-to-UV photon upconversion.82 This two-photon approach is an intriguing development and may open up new reaction possibilities, facilitate photoredox catalysis in aqueous environments, and allow substrates that are challenging to reduce to be accommodated. The downside is that this approach does require high-power laser excitation, so it will likely not become as widespread as traditional photoredox catalysis, where cheap and readily available fluorescent or LED light sources are sufficient. The mer isomers of the Ir(C^N)3 photosensitizer family can be prepared in many cases, but they are isomerized photochemically or thermally to the more stable fac form.83 They can be accessed from dimers of the type Ir1–Ir3 by adding an additional equivalent of the C^N ligand under milder conditions than are needed to form the fac product, and a recently developed strategy allows the mer isomer to be prepared under very mild conditions by transmetalation.84 However, the mer isomers typically have shorter lifetimes than the fac form, and since they are known to photoisomerize they have not been used as photosensitizers.

1.10.5.3 1.10.5.3.1

Cationic bis-cyclometalated iridium complexes Complexes with bipyridine-derived ancillary ligands

Another prominent category of cyclometalated iridium photosensitizers, which has become especially popular in the field of organic photoredox catalysis, has the general formula [Ir(C^N)2(N^N)]+, where N^N is a neutral, chelating diimine ligand from the 2,20 -bipyridine family. These compounds can be prepared by direct reaction of the diimine ligand with chloro-bridged dimers [Ir(C^N)2(m-Cl)]2, e.g. Ir1–Ir3, or in many cases more easily from solvent-bound complexes of the type [Ir(C^N)2(solv)2]+, where the bound solvent is usually MeCN.85 The charge-compensating counterion is most often a noncoordinating ion such as PF–6. Complexes of this type were first introduced and extensively studied by Watts et al.,65,86–88 including the derivatives Ir23–Ir28 shown in Fig. 10. Characterization of the photophysical and redox properties of these early examples revealed that these cationic

Fig. 10 Early examples of [Ir(C^N)2(N^N)]+ complexes prepared by Watts et al.

296

Organometallic Photosensitizers

complexes are weaker photoreductants but stronger photooxidants than fac-Ir(ppy)3 derivatives, but in the early years after their discovery they were not specifically evaluated as photosensitizers. Following this early work that introduced the general [Ir(C^N)2(N^N)]+ structure type, there was steady research that outlined additional derivatives within this class, including versions with substituted and p-extended N^N ligands,89–93 some capable of elaboration into multimetallic structures,94,95 and variants with substituted C^N or N^N ligands suitable for conjugation to biomolecules.96–100 Most of these previous studies included detailed photophysical and electrochemical characterization, but none specifically looked at the function of these compounds as photosensitizers. An explosion in the structural variety of cationic [Ir(C^N)2(N^N)]+ complexes, and the beginnings of their applications as photosensitizers, emerged in large part due to the efforts of Bernhard and co-workers. After originally studying the complex [Ir(ppy)2(dtbbpy)](PF6) (Ir29, dtbbpy ¼ 4,40 -di-tert-butyl-2,20 bipyridine) and showing it to be effective for thin-layer electroluminescence,101 they developed a parallel synthesis approach that allowed rapid synthesis of and evaluation of >100 variants of this general structure type (see Fig. 11).102 Their series includes several of the known compounds shown in Fig. 10, but the combinatorial method they introduced allowed rapid access to many other related complexes and evaluation of their photoluminescence. A subset of the compounds shown in Fig. 11 were evaluated as photosensitizers for the catalytic reduction of protons to hydrogen, in combination with a cobalt-based electrocatalyst,103 whereas others were evaluated in combination with a rhodium-based water reduction catalyst.104 This work demonstrated the potential of these cationic iridium complexes as visible-light photosensitizers for catalysis, with the high throughput approach allowing rapid synthesis and screening of numerous photosensitizer candidates.

Fig. 11 Structures of cationic bis-cyclometalated iridium photosensitizers prepared combinatorially by Bernhard et al.

The combinatorial approach described above allowed many empirical structure-property relationships to be established. A more focused study on a smaller set of cationic bis-cyclometalated iridium photosensitizers provided insights into how the substitution pattern of the C^N ligand can be used to systematically control the ground- and excited-state redox properties. The complex Ir29 (Fig. 11), along with the three substituted variants shown in Fig. 12 (Ir30–Ir32), were compared.105 The sequential addition of electron-withdrawing groups to the cyclometalating ligands had the effects of stabilizing the HOMO, i.e. shifting the formally IrIV/IrIII redox couple to more positive potentials, while also raising the energy of the T1 excited state. The potential of the first reduction event, which is localized on the dtbbpy ancillary ligand, also shifted positively but by a much smaller amount. As a result, complexes Ir29–Ir32 are all similarly potent as photoreductants, but complex Ir32, with the most electron-withdrawing substituent set, is a significantly stronger photooxidant than other cyclometalated iridium photosensitizers. This complex is also a reasonably

Organometallic Photosensitizers

297

strong photoreductant, and in the original report Ir32 was shown to outperform [Ru(bpy)3]2+ as a photosensitizer for visible-light induced proton reduction, in terms of both TON of H2 and quantum yield. Since these original reports, complexes Ir29–Ir32, in particular Ir29 and Ir32, have emerged as extremely popular photosensitizers for applications in photoredox organic synthesis methodology. The first reports of Ir29106 and Ir3271 in organic synthesis applications appeared over 10 years ago, and since then they can be found in hundreds of reports on photoredox methodology,75,107–127 with the cited references representing but a small fraction of those in the literature. Whereas Ir29 and Ir32 have become and remain the two most popular cationic iridium photosensitizers, other closely related variants have been developed, also summarized in Fig. 12. Originally part of a combinatorial library introduced by Bernhard et al.,104 Ir33, with an unsubstituted 2,20 -bipyridine ancillary ligand, has become prominent in organic synthesis applications, particularly in combination with nickel co-catalysts that participate in electron- or energy-transfer steps with the excited iridium photosensitizer.128–130 Derivatives with electron-withdrawing groups on the bpy ancillary ligand (Ir34–Ir36) have also emerged, and these are especially prominent in transformations involving challenging photoinduced oxidation steps.131 The [Ir(C^N)2(N^N)]+ framework is synthetically more simple than homoleptic Ir(C^N)3 analogues, and redox potentials can be readily tuned to allow both reductive and oxidative SET pathways, making these complexes attractive for a wide range of photoredox applications.

Fig. 12 Structures of cationic bis-cyclometalated iridium photosensitizers modified with electron-withdrawing groups.

1.10.5.3.2

Carboxy-substituted complexes for solar cells

Related cationic cyclometalated iridium complexes (Fig. 13) have also been applied as photosensitizers in dye-sensitized solar cells (DSSCs), in which the excited photosensitizer injects electrons into a semiconductor support. One of the earliest reports on cyclometalated iridium complexes in DSSCs was in 2006, where two complexes with carboxylic acid groups on the N^N ligand, Ir37 and Ir38, were evaluated as dyes for TiO2 solar cells.132 These compounds had much less intense visible absorption than the commonly used carboxylated complex [Ru(bpy)2(dcbpy)]2+ (dcbpy ¼ 4,40 -di-carboxy-2,20 -bipyridine), but a steep decrease in lifetime when deposited on TiO2 indicated very efficient charge injection into TiO2. Their overall solar-to-current efficiency was 0.5% (Ir38) or 0.65% (Ir37), only about half that of an analogous cell with the [Ru(bpy)2(dcbpy)]2+ dye. The synthesis and photophysical properties of complex Ir39, a carboxylated derivative of Ir23, were first described by Amouri et al.133 This complex and two analogues, Ir40 and Ir41, were evaluated as dyes for DSSCs,134,135 albeit with poor efficiency, in large part because cyclometalated iridium complexes of this type are poor absorbers over most of the solar spectrum, in contrast to ruthenium polypyridyl complexes which are ubiquitous in DSSCs.136 Complex Ir42, with a modified anchoring group on the bpy ligand, was prepared but no improvements in device performance were apparent.137 Complex Ir39 has also been used to construct coordination polymers and metal-organic frameworks (MOFs), with the carboxylate groups serving as linkers to zinc138 or yttrium139 nodes. The photoreducing ability of the bis-cyclometalated iridium centers allowed the zinc materials to function as luminescent sensors for nitroaromatic explosives, via electron-transfer quenching of the phosphorescence,138 with the yttrium materials being effective for visible-light promoted CO2 reduction.139

298

Organometallic Photosensitizers

Fig. 13 Structures of carboxylic acid-functionalized cationic bis-cyclometalated iridium complexes, used in DSSCs and photoactive coordination polymers.

1.10.5.3.3

Cationic complexes with quinoline-derived cyclometalating ligands

One of the major drawbacks of the [Ir(C^N)2(N^N)]+ photosensitizers described above is their limited absorption of visible light, particularly when C^N ¼ ppy or a substituted analogue. To improve the visible-light absorption of this photosensitizer class, several groups have investigated analogues where the cyclometalating (C^N) ligand and/or the N^N ligand are more highly conjugated. In complexes of this type the HOMO is normally delocalized over the Ir center and the phenyl groups of the C^N ligand, whereas the LUMO is primarily localized on the N^N ligand. As a result, altering the conjugation of the cyclometalating ligand normally has only a small effect on the HOMO energy, but the LUMO can be readily tuned by increasing the conjugation of the N^N ligand. An early systematic study on complexes with increased conjugation in the C^N and N^N ligands disclosed by Li et al. included the complexes summarized in Fig. 14 below. Although these compounds were not specifically targeted as photosensitizers, this work did reveal key structure property relationships that set the stage for further development of cationic bis-cyclometalated iridium photosensitizers with enhanced visible absorption. Whereas complex Ir23 and its closely related derivatives above have peak MLCT absorptions at 400 nm and are completely transparent beyond 500 nm, complexes Ir43–Ir48 in Fig. 14 all have peak MLCT absorption wavelengths 440 nm, with significant absorption tailing well beyond 500 nm.

Organometallic Photosensitizers

299

Fig. 14 Structures of cationic bis-cyclometalated iridium complexes with enhanced conjugation and increased visible-light absorption.

In addition to the piq cyclometalating ligand used in many of the above examples, the isomeric cyclometalating ligand 2-phenylquinoline (pq), used in Ir41 above, also has the effect of increasing the visible absorption. A recent study outlined a series of cationic complexes with pq cyclometalating ligands, where the ancillary ligand(s) were either monodentate substituted pyridines or substituted 2,20 -bipyridines (Fig. 15).140 This work found that complexes with C^N ¼ pq typically have longer lifetimes than reference compound Ir23, in aerated solutions. Moreover, comparing complexes with monodentate pyridine ancillary ligands (Ir49–Ir52) to those with substituted bipyridine (Ir53–Ir56), the latter are more difficult to oxidize and easier to reduce, on account of the stronger-field ancillary ligand stabilizing the t2g HOMO, and the accessible LUMO of the bpy derivative. Photosensitization of singlet oxygen was studied in this work, which found that bidentate complexes Ir54–Ir56 typically had higher quantum yields for singlet oxygen generation, with complex Ir56 performing the best. Relatedly, these compounds are also effective photocatalysts for oxidation of sulfides to sulfoxides, which proceeds via singlet oxygen.

Fig. 15 Structures of Ir photosensitizers used for singlet oxygen generation and photocatalytic sulfide oxidation.

300

Organometallic Photosensitizers

1.10.5.3.4

Cationic bis-cyclometalated iridium photocatalytic dyads

Much of the early work on the [Ir(C^N)2(N^N)]+ structure class examined these complexes as photosensitizers for photocatalytic hydrogen evolution, and in more recent years complexes related to those described above have been used as photosensitizers for hydrogen evolution. In addition, dyad structures have been constructed, shown in Fig. 16, where the iridium photosensitizer is linked to the hydrogen evolution catalyst. As an example, dyad Ir57 was prepared, which links a cationic iridium photosensitizer to a cobalt hydrogen evolution electrocatalyst. This complex was able to catalyze hydrogen evolution under visible-light irradiation, in acetone solution with NEt3 and HNEt3BF4 present as the electron donor and proton source, respectively.141 Dyad Ir57 performed similarly to systems where Ir26 was mixed with [Co(dmgBF2)2(OH2)2] under otherwise identical conditions. In an effort to red-shift the absorption window to capture more visible light, Hanan, Elias, prepared the dyad Ir58, which uses the piq cyclometalating ligand known to increase visible absorption.142 In their complexes the iridium photosensitizer was joined with Co(dmgH)2(Cl), closely related to the cobalt catalyst used in Ir57. In this work they showed that visible light excitation in MeCN solution with HBF4 as the acid and triethanolamine (TEOA) as the base produced hydrogen, and found that dyad Ir58 substantially outperformed bimolecular systems where Ir43 was combined with the Co catalyst. In more recent work on related systems, they prepared an expanded series of complexes, varying the position of the pyridine linker used to coordinate the cobalt catalyst, and also

Fig. 16 Structures of cyclometalated iridium-cobalt dyads used for photocatalytic hydrogen evolution.

Organometallic Photosensitizers

301

altering the cyclometalating ligand on iridium.143 In this work they found that all of dyads Ir58–Ir63 outperformed bimolecular combinations of photosensitizer and catalyst, in terms of turnover frequency (TOF) and turnover number (TON) for hydrogen evolution promoted by visible light. Photosensitizers related to those shown above in Fig. 14 were also used for photocatalytic CO2 reduction. The complex Ir64, a methyl-substituted analogue of Ir43 was prepared as the photosensitizer, and a related dyad Ir65 was also accessed where the photosensitizer is joined with a rhenium carbonyl CO2 reduction catalyst (Fig. 17). Methyl substitution in Ir64 makes the complex 90 mV more difficult to reduce, but otherwise the electrochemical and photophysical properties are virtually identical to unsubstituted Ir43. Ir64 and Re(dmb)(CO)3Br (dmb ¼ 4,40 -dimethyl-2,20 -bipyridine) or Ir65 by itself are both active for carbon dioxide reduction under visible light irradiation. Typical reaction conditions involved the catalyst system in the presence of CO2 with filtered white light irradiation (l > 500 nm), DMA/TEOA (5:1) as the solvent, and a nicotinamide or dihydroimidazole sacrificial reductant. These systems reduce CO2 to CO selectively, with a small amount of H2 detected as well. Again, the enhanced visible-light absorption brought on by the piq cyclometalating ligand was crucial for photocatalytic performance with visible light, though in this system the dyads did not dramatically outperform mixtures of Ir64 and the rhenium catalyst, in terms of TON of CO production and selectivity for CO over H2 and HCOOH.

Fig. 17 Structures of Ir photosensitizer and Ir-Re dyad used for photocatalytic CO2 reduction.

1.10.5.3.5

Complexes with alternative diimine ancillary ligands: enhanced light absorption and charge separation

The majority of the complexes described above feature ancillary ligands derived from 2,20 -bipyridine, 1,10-phenanthroline, or a closely related substituted derivative. There have been some reports of cationic [Ir(C^N)2(N^N)]+ complexes where the N^N ligand is not in the bipyridine family. A thorough study reported in 2011 described 13 photosensitizers with varied N^N ligands, summarized in Fig. 18.144 This study includes seven complexes, Ir66–Ir72, where the N^N ligand is still a substituted 2,20 -bipyridine, but with the

Fig. 18 Structures of iridium photosensitizers with modified N^N ligands, evaluated for photocatalytic hydrogen evolution.

302

Organometallic Photosensitizers

substituents on the 6 and/or 60 positions of the bpy ligand, a substitution pattern not previously studied in the development of iridium photosensitizers. In addition, alternative asymmetric N^N ligands where pyridine is paired with other cyclic or acyclic N-donor moieties (Ir73–Ir77) were also prepared. The bpy substituents in Ir66–Ir72 influence the redox potentials and PL wavelengths in a predictable way, with electron-withdrawing groups anodically shifting the first reduction to more positive potentials and red-shifting the PL spectrum. Complexes with alternative N^N ligands beyond the bpy family typically have very weak or no photoluminescence, with reduction potentials that are either more negative (Ir75 and Ir76) or more positive (Ir73, Ir74, and Ir77) than parent bpy complex Ir23. The UV-vis absorption profiles for all complexes in this series are all quite similar, although Ir74 does have a broad, weak band that tails further into the visible range (l < 600 nm) than the rest. These complexes were evaluated as photosensitizers for visible-light promoted hydrogen evolution, in combination with an iron carbonyl water-reduction catalyst. Only the complexes with substituted bpy ancillary ligands were active for hydrogen evolution, with Ir66 and Ir67 giving the highest turnover numbers. More dramatic perturbation of the electronic structure can be realized by incorporating additional nitrogen atoms into the chelating heterocyclic core, in particular replacing pyridine rings with pyrazine or other diazines. Since the LUMO in these cationic complexes is localized on the N^N ligand, and since the ligand-field strength of the N^N ligand affects the HOMO energy, these modifications can have measurable effects on both frontier orbitals. Such design strategies are replete in the field of ruthenium polypyridyl photosensitizers,145 and have more recently been explored in the design of cyclometalated iridium analogues. Four cationic [Ir(C^N)2(N^N)]+ complexes with systematically varied N^N ligands are shown in Fig. 19. These complexes were designed to be strong photooxidants, and were evaluated for their ability to photodamage electron-rich guanine biomolecules.146 Compared to reference compound [Ir(ppy)2(bpy)]+ (Ir23), complexes Ir78–Ir81 are all >700 mV more difficult to oxidize, with the formally IrIV/IrIII couple occurring outside the MeCN solvent window by virtue of the strong electron-withdrawing power of the CF3 groups. The effect of the N^N ligand is clearly seen by comparing the one-electron reduction potentials, which involve population of the N^N LUMO. Complex Ir78 is easier to reduce by 230 mV compared to unsubstituted Ir23, indicating that there is some electronic coupling of the C^N and N^N ligands, whereby modification of the C^N ligand can influence the N^N LUMO. More dramatic effects are revealed in Ir79–Ir81, where increasing nitrogen content in the N^N ligand leads to progressively more positive reduction potentials. At the extreme, complexes with 2,20 -bipyrazine (Ir80) and 1,4,5,8,-tetraazaphenanthrene (Ir81) as the ancillary ligand are ca. 500 mV easier to reduce than the bpy analogue Ir78. Additional nitrogen atoms in the N^N ligand do not dramatically alter the UV-vis absorption profile, but they do progressively red-shift the photoluminescence, by as much as 86 nm (3000 cm−1) in complex Ir80 compared to Ir78. In terms of their photoredox properties, complexes Ir80 and Ir81 are the strongest photooxidants in the series, and their photoinduced electron-transfer reactions with the electron donors 1,4-hydroquinone and deoxyguanosine50 -monophosphate (dGMP) were faster than the other examined complexes, as determined by Stern-Volmer quenching studies.

Fig. 19 Bis-cyclometalated iridium photooxidants with variably conjugated N^N ligands.

Organometallic Photosensitizers

303

Through a combination of the same electron-deficient C^N ligands and more highly conjugated N^N ligands, it is possible to prepare [Ir(C^N)2(N^N)]+ complexes that retain the strong photooxidizing properties of Ir81 but are even stronger absorbers of visible light. Complexes Ir82 and Ir83 are p-extended analogues of Ir81, and these three complexes were prepared and their photoredox and light-absorption properties were compared.147 Complexes Ir81–Ir83 are equally strong as photooxidants, with excited-state potentials E( IrIII/IrII) that are within 20 mV of each other but all 200 mV more positive than that of commonly used photooxidant Ir32. Although the excited-state redox properties of these three compounds are nearly identical, the visible absorption sequentially improves as the conjugation is increased, as judged by the molar absorptivity (e) at 450 nm. The e value at 450 nm increases from 500 M−1cm−1 (Ir81) to 1400 M−1cm−1 (Ir82) to 9800 M−1cm−1 (Ir83) as conjugation is added to the N^N ligand. The visible absorption profile of Ir83 is almost identical to the complex [Ru(bpy)3]2+, which has been prolifically applied as a visible-light photosensitizer.53 All of Ir81–Ir83 were shown to be effective photooxidants capable of efficient excited-state electron transfer with donor substrates, including organic substrates as well as the halide quenchers chloride, bromide, and iodide. The complexes all oxidized the quenchers with similar rate constants following visible-light excitation, with a Marcus theory analysis of the results indicating small reorganization energies for the rapid electron-transfer events. As described above, a major drawback of most [Ir(C^N)2(N^N)]+ photosensitizers is that they are poor absorbers of visible light, especially in comparison to the ubiquitous ruthenium polypyridyl family of complexes which has been very successful in solar cell and other light-harvesting applications.14 Some of the quinoline-based complexes described above in Figs. 14 and 15 exhibit subtle enhancements in their UV-vis absorption profile compared to complexes where C^N is a phenylpyridine derivative, and as just mentioned above the complex Ir83 is a strong light absorber on account of the more conjugated N^N ligand. But further improvements in the light-harvesting attributes of cyclometalated iridium complexes have been realized by more dramatic structural changes to the C^N and N^N ligands. An early breakthrough in the design of panchromic iridium complexes came from Hasan and Zysman-Colman, who moved away from the typical bipyridine-derived ancillary ligand framework and used bis[(4-methoxyphenyl)imino]acenaphthene (BIAN) ligands to improve light harvesting.148 Both of complexes Ir84 and Ir85 (Fig. 20) exhibited higher molar absorptivities throughout the UV-visible region, as compared to reference compound Ir23. In Ir84 the broad visible absorption covered the entire visible spectrum out to 800 nm, and its light-harvesting efficiency is as high as the popular panchromic ruthenium dye N3.149 However, the excited-state potentials E(IrIV/ IrIII) for Ir84 and Ir85 were not sufficiently negative to allow electron injection into TiO2, so these complexes were not evaluated as solar cell dyes. In subsequent work, the same authors explored the effects of the BIAN aryl substituents on the light-harvesting properties, in a study that included Ir85–Ir90.150 This work showed that Ir86, where the BIAN aryl substituent is NMe2, has the smallest electrochemical HOMO–LUMO gap in the series, as well as the lowest excited-state E0,0 value. Complex Ir86 also has the most intense visible absorption in the series, with a strong band at 576 nm (e ¼ 1.42  104 M−1cm−1) that is assigned as a mixed MLCT/LLCT transition. In the most recent work on this series of compounds, Ir91–Ir94 were prepared, which involve substitution on the C^N ligand as well as at the bay region of the BIAN core.151 All of Ir91–Ir94 are very similar to Ir86 in terms of electrochemical properties, HOMO–LUMO gaps, and UV-vis absorption profiles. This indicates that the substitution pattern of the C^N ligand and peripheral substitution of the N^N ligand have very minor effects on orbital energies and excited states, as has been established for many other classes of cationic bis-cyclometalated iridium complexes.

Fig. 20 Structures of panchromic bis-cyclometalated iridium complexes with BIAN ancillary ligands.

Derivatives of Ir85 and Ir87 with carboxylic acid substituents on the C^N ligand, as well as other complexes with more typical N^N ligands were evaluated as dyes in p-type solar cells.152 These complexes are summarized in Fig. 21 below, and again the BIAN complexes (Ir100 and Ir101) had the most intense visible light absorption. Solar cells were fabricated with all of Ir95–Ir100, and while the efficiencies were modest for all complexes, the panchromic-absorbing Ir100 did have the highest efficiency.

304

Organometallic Photosensitizers

Fig. 21 Structures of carboxylic acid-substituted bis-cyclometalated iridium photosensitizers used in solar cell design.

In addition to carboxylic acid-decorated C^N ligands, phosphonate analogues have also been investigated in the construction of solar cell dyes. Fig. 22 summarizes three dyes that were evaluated in p-type DSSCs.153 These complexes were designed to anchor the C^N ligands to the NiO semiconductor, such that following hole injection the remaining electron would be delocalized onto the N^N ligand, ensuring a long-lived charge separation with slow recombination. The increased conjugation in Ir103 and Ir104 augmented the molar absorptivity in the near-UV and blue regions, relative to 1,10-phenanthroline complex Ir102, but there was no marked improvement in absorption throughout the rest of the visible range. Excited-state lifetimes for these complexes are short in

Fig. 22 Structures of phosphonate-substituted bis-cyclometalated iridium photosensitizers used in p-type DSSC design.

Organometallic Photosensitizers

305

solution, 60 ns, but when grafted to TiO2 transient absorption signatures for the reduced iridium complexes are observed, with the charge-separated state persisting on the microsecond timescale. The efficiencies of p-type solar cells constructed from these complexes were still modest, on account of the poor visible light absorption, but this work does provide important insight into controlling charge separation and injection. Another strategy to obtain panchromic visible absorption in cyclometalated iridium complexes is to use azoimidazole ancillary ligands, which can coordinate in two different protonation states and support neutral (Ir105) or cationic (Ir106) bis-cyclometalated iridium complexes, shown in Fig. 23.154 Protonation of the ancillary ligand in Ir106 stabilizes both the HOMO and the LUMO, the latter by a substantially larger amount, leading to a HOMO–LUMO gap that is smaller in Ir106 by almost 0.4 eV. In these compounds the HOMO and LUMO are both localized on the ancillary ligand. Both complexes absorb strongly in the visible region, with the peak absorption in Ir106 (l ¼ 565 nm, e ¼ 3.38  104 M−1 cm−1) red-shifted by 2850 cm−1 and more intense than that of neutral complex Ir105 (l ¼ 493 nm, e ¼ 2.21  104 M−1 cm−1). OPV devices were constructed with Ir106 as the acceptor material, and while the device had a high open-circuit voltage of 1.12 V, the current values were very low. Despite some of the design improvements highlighted above, which have led to improved solar capture for cationic cyclometalated iridium complexes, these complexes still fall short of the well-known ruthenium polypyridyl dyes as solar light harvesters, in particular for photovoltaic applications. The iridium complexes that have been tested in solar cells are either poor absorbers of visible light, or if they do absorb well they have short excited-state lifetimes and/or redox potentials that are not conducive to semiconductor charge injection. There has yet to be a breakthrough to make iridium dyes competitive with ruthenium dyes for DSSCs, and it is not immediately clear which design features would lead to such a breakthrough.

Fig. 23 Structures of proton-switchable panchromic absorbers with azoimidazolate ancillary ligands.

1.10.5.3.6

Alternative cyclometalating ligands for enhanced light absorption and charge-transfer lifetimes

The examples provided above highlighted how changes to the N^N ligand can bring about significant changes in the UV-vis absorption profile and excited-state lifetimes. There are some examples above (e.g. Ir78–Ir83) that also used substituents on the C^N ligand to perturb frontier orbital energies, especially the HOMO, which likewise has an effect on the excited-state redox properties. Moving beyond 2-phenylpyridine and phenyl-substituted quinoline derivatives, there are other classes of cyclometalating ligands that can have more dramatic effects on the excited states of [Ir(C^N)2(N^N)]+ complexes. One cyclometalating ligand class which has been used extensively in the design of phosphorescent iridium complexes is the substituted 2-phenylbenzothiazole series, often abbreviated as “bt” or “pbt”. Yu et al. prepared a series of bis-cyclometalated iridium complexes with substituted bt cyclometalating ligands and 4,40 -dicarboxy-2,20 -bipyridine (dcbpy) as the N^N ligand.155 In this case the complexes were isolated with dcbpy in the monoprotonated state, such that the complexes were overall charge neutral, as shown in Fig. 24. Complexes Ir107–Ir110, either unsubstituted or with CF3 groups at different positions on the bt ligands, all have very similar UV-vis absorption profiles, with slightly more visible absorption than complexes with 2-phenylpyridine-derived C^N ligands but still only absorbing blue light in the visible region with e 1  104 M−1 cm−1 for their MLCT bands. In contrast, dimethylamino-substituted Ir111 has much stronger visible absorption, with e 4  104 M−1 cm−1 but still with absorption only in the blue region. This much more

Fig. 24 Structures of bis-cyclometalated iridium photosensitizers with substituted 2-phenylbenzothiazole cyclometalating ligands.

306

Organometallic Photosensitizers

intense band in Ir111 was proposed to arise from an intraligand (IL) transition on the C^N ligand, likely due to the “push-pull” electronic nature of this ligand. In these complexes the one-electron reduction potential is quite sensitive to the substituent pattern on the C^N ligand, even though DFT indicates that the LUMO remains localized on the N^N ligand. These complexes were evaluated as photosensitizers for visible-light promoted hydrogen evolution, in combination with colloidal Pt in aqueous solution at pH 7. Despite its strong absorption Ir111 is minimally active for photocatalytic hydrogen evolution, with Ir107 and Ir108 performing the best, the latter giving 1.5  103 turnovers after 20 h of irradiation. These complexes were also adsorbed onto TiO2, and hydrogen evolution was studied under otherwise very similar conditions. Again, Ir107 and Ir108 outperformed the other complexes, although in these experiments Ir111 did not perform as poorly. Finally, all five complexes were tested as photosensitizers in DSSCs, all giving open-circuit voltages (VOC) of ca. 0.5 V and short-circuit current densities that ranged from 2.2 to 3.7 mA/cm2. Complex Ir109 resulted in the highest solar cell efficiency, 1.39%, which is still lower than ruthenium-based DSSCs but better than the other iridium complexes referenced above. The cyclometalating ligand 2-pyridyl-benzothiophene, often abbreviated as btp, is an isomer of bt (see Fig. 24) that has likewise been used in cationic [Ir(C^N)2(N^N)]+ designs. A series of complexes with variably substituted bpy ligands and btp as the cyclometalating ligand was described by Takizawa et al.,156 and these are summarized in Fig. 25. Complexes Ir112–Ir115 all have 1MLCT absorption transitions in the visible region, occurring at very similar wavelengths (428–437 nm) and with very similar molar absorptivities as the bt complexes summarized above in Fig. 24. In these complexes the MLCT transition involves the btp cyclometalating ligand, not the bpy, such that electron-withdrawing groups on the bpy blue-shift the absorption band, with the opposite observed for electron-donating groups. Complexes Ir112–Ir115 were evaluated as photosensitizers for electron-transfer reactions with TEOA and methyl viologen (MV2+). No quenching was observed with TEOA, suggesting these compounds are not active as photooxidants, but Ir112–Ir114 did exhibit fast electron transfer to MV2+, indicating they are potent photoreductants. Quenching was most efficient with Ir114, and it was suggested by the authors that the stronger s-donation in bpy ligands substituted with electron-donating groups destabilizes the higher-lying 3MC states, improving the lifetime of the 3MLCT state and resulting in more efficient bimolecular quenching.

Fig. 25 Structures of bis-cyclometalated iridium photosensitizers with substituted 2-pyridylbenzothiophene cyclometalating ligands.

Another cyclometalating ligand class which has been used in the design of iridium photosensitizers derives from coumarins, a well-known class of organic fluorophores. Takizawa et al. introduced cationic bis-cyclometalated iridium photosensitizers with bipyridine-derived ancillary ligands and used them for visible-light-driven hydrogen generation.157 The structures of these complexes are summarized in Fig. 26. One advantage of this design is the enhanced visible absorption brought on by the coumarin ligand, with e > 1.2  105 M−1 cm−1 for Ir116 and Ir117. In addition, these compounds have long-lived excited states, with room-temperature lifetimes between 6 and 7 ms for Ir116 and Ir117 (Ir118 is not luminescent). Finally, all of Ir116–Ir118 were tested as photosensitizers for light-driven hydrogen evolution, using [Co(bpy)3](Cl)2 as the catalyst and triethylamine as the sacrificial reductant. Complexes Ir116 and Ir117 both outperformed Ir23 in this reaction, with TON 1500 relative to photosensitizer loading, about five times more effective than Ir118 and Ir23. In a subsequent study, the same group investigated substituent effects in the same class of complexes, as well as a complex with a diamine instead of a bipyridine ancillary ligand (Ir123).158 These variants are also summarized in Fig. 26. The UV-vis absorption and emission spectra of all complexes were quite similar, although trifluoromethyl-substituted complex Ir122 was not luminescent in either solvent tested (CH2Cl2 and CH3CN) and Ir116 did not luminesce in CH3CN. The photoluminescence quantum yield depended both on the ancillary ligand and the solvent, with diamine complex Ir123 having the highest photoluminescence quantum yield in both CH2Cl2 (FPL ¼ 0.583) and CH3CN (FPL ¼ 0.429). These complexes were also evaluated as photosensitizers for visible-light driven hydrogen evolution, this time with a cobalt glyoxime co-catalyst and ascorbate as the sacrificial reductant. Complexes with electron-donating groups on the bipyridine (Ir119–Ir121) vastly outperformed the others. Mechanistic studies revealed that the excited-state of complex Ir121 was much more efficiently quenched by ascorbate than that of Ir116, suggesting that the origin of the superior photocatalytic activity in the latter is the more rapid reaction with the sacrificial reductant.

Organometallic Photosensitizers

307

Fig. 26 Structures of coumarin-based cationic bis-cyclometalated iridium photosensitizers.

Castellano et al. have shown that bis-cyclometalated iridium complexes can be prepared using 1,8-naphthalenebenzimidizole (NBI), another well-studied organic chromophore, as the cyclometalating ligand (Fig. 27). Complex Ir124 and its homoleptic facIr(C^N)3 analogue were prepared and shown to have efficient red phosphorescence with lifetimes > 1 ms.159 In subsequent work, the same group showed that Ir124 is a robust visible-light photosensitizer for photocatalytic hydrogen evolution, in combination with [Co(dmgH)2(pyridine)(Cl)] as the proton reduction catalyst and N,N-dimethyl-p-toluidine (DMT) as the sacrificial reductant.160 Complex Ir124 has multiple intense visible absorption bands centered at 452 nm attributed to 1MLCT transitions, and an excited-state redox potential E(IrIV/ IrIII) of −0.95 V vs Fc+/Fc that is less reducing than many of the cationic photosensitizers described above (e. g. Ir29 and Ir32) but still sufficient for proton reduction. Prolonged irradiation of Ir124, [Co(dmgH)2(pyridine) (Cl)], and DMT in 2:1 MeCN/H2O leads to sustained hydrogen production that levels off after 20 h but continues when more cobalt catalyst and DMT are added. After 90 h of irradiation there was no evidence of any decomposition of Ir124, suggesting it is a robust photosensitizer under these conditions and capable of prolonged operation. In a direct comparison under identical

Fig. 27 Cationic heteroleptic complexes with 1,8-naphthalenebenzimidizole and pyrene-derived cyclometalating ligands.

308

Organometallic Photosensitizers

conditions, Ir124 had both a faster initial rate of H2 evolution and 5 as many turnovers as [Ru(bpy)3]2+ during the first 20 h of irradiation. Pyrene-based cyclometalating ligands have also been developed as a strategy to increase lifetimes of [Ir(C^N)2(N^N)]+ photosensitizers. Pope et al. prepared the complexes summarized in Fig. 27 which partner cyclometalating pyrene ligands with bpy ancillary ligands.161 The visible absorption of these compounds is primarily in the blue range, but in most cases with higher molar absorptivities than the reference compound Ir23. The excited-state lifetimes are 13.3 ms (Ir125) and 3.9 ms (Ir126), significantly longer than Ir23 measured under the same conditions (0.34 ms). The complexes are efficiently quenched by 3O2, producing 1 O2 with near-unity quantum yields, and they were applied as photooxidants for the conversion of 1,5-dihydroxynaphthalene to Juglone.

1.10.5.4 1.10.5.4.1

Charge-neutral heteroleptic bis-cyclometalated iridium complexes Complexes with acac ancillary ligands

Charge-neutral bis-cyclometalated iridium complexes, where the ancillary ligand is normally a monoanionic chelating ligand, are arguably the most widely studied class of cyclometalated iridium complexes. Of these, the most prominent structure type is Ir(C^N)2(acac), where “acac” is acetylacetonate or a closely related analogue. Such compounds have been especially important in the design of organic light-emitting diodes (OLEDs), given their relative ease of synthesis, facile color tunability, and high photoluminescence quantum yields over a large part of the visible spectrum.22,24,162,163 By and large the photophysical and electrochemical properties of Ir(C^N)2(acac) complexes closely resemble the homoleptic fac-Ir(C^N)3 complexes described above. The photophysics in these complexes is largely dominated by mixed 3MLCT/3LC states involving the cyclometalating ligands, and the acac is normally photophysically innocent. This stands in contrast to most of the cationic [Ir(C^N)2(N^N)]+ complexes described above, where the LUMO is often localized on the N^N ligand, and thus the N^N ligand is a critical determinant of the excited state properties. Although these Ir(C^N)2(acac) compounds have mainly been studied as phosphors for optoelectronic devices, some of the earliest work on this class of compounds recognized their potential as photosensitizers. Thompson et al., originators of the Ir(C^N)2(acac) structure class, studied the series of compounds in Fig. 28 as singlet oxygen sensitizers.164 In addition to complexes

Fig. 28 Structures of neutral heteroleptic bis-cyclometalated iridium complexes examined as photosensitizers for singlet oxygen generation.

Organometallic Photosensitizers

309

Ir127–Ir130 with acac or dipivaloylmethane ancillary ligands, glycinate complex Ir131 was investigated, along with Ir132 which has two monodentate ancillary ligands. All of these complexes have high quantum yields for 1O2 generation following excitation with 355-nm (FD ¼ 0.54–0.95) or 532-nm (FD ¼ 0.77–1.00) excitation. Another notable outcome is that quenching of the excited state by 1O2 is quite slow, which allows productive sensitization of 1O2. There was no systematic dependence of the quenching behavior on the ancillary ligands, not surprising since the triplet excited states in iridium complexes of this type are dominated by the cyclometalating ligands. In a later study by some of the same authors, a larger subset of Ir(C^N)2(acac) complexes were studied as singlet oxygen photosensitizers.165 These additional complexes are also summarized in Fig. 28, and in this study the iridium complexes were studied alongside Pt(C^N)(acac) complexes featuring some of the same C^N ligands. The authors measured high singlet oxygen quantum yields for all of the compounds, but the Pt(II) analogues were all near unity whereas the iridium complexes ranged between 0.59 and 0.90. The results of Stern-Volmer quenching studies suggest that iridium complexes form 1O2 via a combination of energy and electron-transfer mechanisms, and that iridium complexes with lower excited-state energies were less efficient at forming singlet oxygen. The authors propose that the planar geometry of the Pt(II) complexes allows more frequent productive collisions between the complex’s excited-state spin density complex and 3O2, which is why the singlet oxygen quantum yields in the platinum complexes tend to be higher. Ir(C^N)2(acac) derivatives and other charge-neutral heteroleptic cyclometalated iridium complexes have been used less frequently in photoredox applications, in comparison to the homoleptic Ir(C^N)3 and cationic [Ir(C^N)2(N^N)]+ complexes described above. Typically for a given cyclometalating ligand the Ir(C^N)2(acac) complexes are weaker photoreductants than fac-Ir(C^N)3 and weaker photooxidants than [Ir(C^N)2(N^N)]+, which at least partially explains their relative scarcity in photoredox methodology. Nonetheless, they are often easier to prepare than homoleptic fac-Ir(C^N)3 complexes and can enable some of the same transformations. Lalavée et al. have studied a wide variety of Ir(C^N)2(acac) derivatives as photoinitiators for polymerization reactions.166,167 The full set of complexes they have examined are shown in Fig. 29. In these reactions the excited iridium complex undergoes a one-electron redox transformation with the initiator. Common initiators include diphenyliodonium salts, reduced directly by the iridium photoinitiator, or aryl bromides which are reduced following reductive quenching of the iridium complex’s

Fig. 29 Structures of Ir(C^N)2(acac) complexes examined as photoinitiators for polymerization reactions.

310

Organometallic Photosensitizers

excited state by an amine. The initial work focused on three fluorinated derivatives where the C^N ligand is a 1-phenylisoquinoline (piq) derivative, Ir137–Ir139. Fluorination in Ir138 and Ir139 resulted in a blue-shifted absorption window, but all three complexes have reasonably strong visible absorption out to at least 550 nm. These complexes were used as light-absorbers for photopolymerization reactions of acrylate and epoxide monomers for radical and cationic polymerization, respectively. In radical polymerization all three of these complexes vastly outperformed fac-Ir(ppy)3 when using 457 nm laser excitation, and this difference was attributed to the superior light absorption by these compounds. For both radical and cationic photopolymerization there was no notable difference in activity between Ir137 and the fluorinated derivatives Ir138 and Ir139. In a more extensive study that included additional fluorinated derivatives Ir140–Ir142, as well as several other acac complexes where the C^N and/or acac ligands are altered, the authors found some variation in activity for photopolymerization reactions. As is normally the case for cyclometalated iridium complexes the UV-vis absorption spectrum is primarily dependent on the C^N ligand, and at the extremes in this series are complex Ir149, which absorbs very weakly beyond 400 nm, and all of the piq-derived complexes (Ir137–Ir142, Ir148, and Ir149) which absorb strongly in the blue and green regions and have significant absorption out to 600 nm. In complexes Ir137, Ir149, and Ir150 the C^N ligand is the same and the substituents on the acac backbone are altered. Comparing these three compounds, the absorption spectra, IrIV/IrIII redox potentials, and triplet-state energies are nearly identical, indicating that the acac plays a very minor role in the frontier orbitals and excited states. In terms of photopolymerization activity, there were no apparent systematic trends as a function of photosensitizer structure, but all tested photosensitizers performed well in light-initiated cationic and radical polymerizations. More recently, Zheng et al. have discovered a strategy for visible-light mediation of the well-known azide-alkyne Huisgen cycloaddition reaction,168 normally catalyzed in the dark by Cu(I). In this work a number of photosensitizers were examined, including the popular candidates [Ru(bpy)3](Cl)2, [Ir(F2ppy)2(dtbbpy)]+ (closely related to Ir32 and prepared as part of the combinatorial library in Fig. 11), and fac-Ir(ppy)3 (Ir4). None of these popular photosensitizers were effective, but very high yields of 1,2,3-triazole products were formed from the azide and alkyne precursors when Ir(piq)2(acac) (Ir151, Fig. 30) was used as the photosensitizer with white LED irradiation. The reaction conditions tolerated a wide range of alkyl azides and both aryl and alkyl terminal alkynes. The reaction is proposed to proceed via initial oxidation of the alkyne by the photosensitizer, which then undergoes addition and cyclization with the azide. The reduced photosensitizer transfers an electron back to the triazole radical cation that forms following cyclization; as a result, the overall reaction is redox neutral and does not require a sacrificial redox reagent. The authors did not propose an explanation why Ir151 was an effective photosensitizer for this method whereas others failed.

Fig. 30 Structure of Ir(piq)2(acac) (Ir151).

Ir(C^N)2(acac) complexes have also been applied as light absorbers in solar cells. In one study carboxylate-substituted derivatives were used as dyes for dye-sensitized solar cells (DSSCs).169 The three complexes studied are summarized in Fig. 31. The presence of the carboxylate groups shifts the PL much deeper into the visible region, in particular for Ir152 where the PL is shifted into the red region, lmax ¼ 634 nm, over 100 nm (3400 cm−1) red shifted from the photoluminescence of unsubstituted Ir(ppy)2(acac) (Ir133). The visible absorption profile in these complexes was typical for cyclometalated iridium complexes, strong in the UV and blue, but weaker in the visible region and tailing to near-baseline by l ¼ 600 nm. Complex Ir154 was the strongest visible-light absorber in the series, on account of the increased conjugation on the C^N ligand. TiO2 solar cells were fabricated with these three dyes, either on transparent TiO2 or transparent TiO2 with a second scattering layer to increase light harvesting. With these three dyes an inverse relationship between short-circuit current density (JSC) and open-circuit voltage (VOC) was noted. Complex Ir152 was at one extreme with the lowest JSC and highest VOC, whereas Ir154 was at the other extreme with the highest JSC and the lowest VOC. The addition of the second TiO2 scattering layer had minimal effect on the VOC but did lead to larger JSC values. All solar cells tested had similar PCE values ranging between 1.96% and 2.51%, and the authors ascribed the modest performance to the modest light-harvesting of these dyes.

Organometallic Photosensitizers

311

Fig. 31 Structures of carboxylate-substituted Ir(C^N)2(acac) complexes used as dyes for DSSCs.

Because of their high solubility and relatively low polarity, Ir(C^N)2(acac) complexes can be blended or layered with common solar cell materials either via solution-processed routes for by thermal evaporation techniques. In one of the earliest reports on this strategy, Yu et al. inserted a 4-nm layer of Ir155 (Fig. 32) into a layered organic solar cell device consisting of bathophenanthroline, C60, and pentacene layers; Ir155 was deposited in between the C60 and pentacene layers.170,171 Ir155 is a tert-butyl-substituted analogue of Ir130, which presumably was chosen to facilitate deposition. In these devices the pentacene layer served as the primary visible light absorber, although the C60 and Ir155 did contribute to absorption in the blue region. The device doped with Ir155 had a higher open circuit voltage than the undoped solar cell, resulting in a higher power conversion efficiency. The long triplet exciton lifetime in the Ir layer allows this compound to function well in a solar cell and permits efficient carrier diffusion, and the proposed benefit of the iridium complex is that its HOMO energy level is lower than that of pentacene, which allows for higher open circuit voltage. The effect of iridium complex doping on a simpler device architecture was also investigated.172 In this work the devices

Fig. 32 Structures of Ir(C^N)2(acac) complexes used in solar cell design.

312

Organometallic Photosensitizers

consisted of a copper phthalocyanine donor layer with a C60 acceptor layer; the C60 was either undoped, doped with Ir156 (Fig. 32), or doped with a red organic fluorophore. Ir156 uses the same coumarin-based C^N ligand found in Ir116–Ir123 described above, this time in an Ir(C^N)2(acac) structure. Devices doped with Ir156 had the best power conversion efficiencies, in this case primarily due to an increase in current density, with nearly identical open circuit voltages for the three devices. The authors propose that the long triplet-state lifetime of Ir156 enables better carrier diffusion in the solar cell. The authors also proposed that Ir156 contributed to higher solar absorption, although the absorption of the Ir-doped cell was only minimally higher than that of the undoped cell, and only in the blue region. Nevertheless, this work likewise demonstrated possible improvements in organic solar cells doped with iridium photosensitizers. Solution-processed solar cells doped with Ir(C^N)2(acac) complexes have also been studied. Zhen et al. introduced two new bis-cyclometalated iridium complexes as electron donors for solution-processed C71 fullerene solar cells.173 Whereas unsubstituted complex Ir144 is reported to be poorly miscible with the fullerene acceptors, substituted complexes Ir157 and Ir158 are modified at the pyridyl ring and they permit fabrication of solution-processed bulk heterojunction solar cells. These complexes were prepared by on-complex Suzuki or Stille coupling reactions with the bromo-substituted precursor. These two complexes have HOMO energy levels that are within 0.05 eV of one another and unsubstituted Ir144. The LUMOs reside on the C^N ligand and that of Ir158 was measured to be 0.30 eV more stable than that of Ir144 on account of the increased conjugation. The LUMO energy of Ir157 was not precisely determined but was estimated to be intermediate between those two values. Both complexes Ir157 and Ir158 have reasonably strong MLCT absorption bands 500 nm, and their absorption throughout the UV-visible region is strongly overlapped with that of the fullerene. The photon to current efficiency (PCE) depended slightly on the Ir/fullerene ratio, and was generally higher for Ir158 compared to Ir157, maximizing at 2.0%. The excited-state lifetimes of the compounds were measured to be 108 ns (Ir157) and 122 ns (Ir158) in neat films, longer than a typical organic material but still somewhat short for a phosphorescent dopant. This work does show that cyclometalated iridium complexes can have appropriately placed HOMO and LUMO energy levels to function in bulk heterojunction solar cells, although overall PCE remained modest. In an attempt to increase the visible absorption of Ir(C^N)2(acac) complexes for solar cells, Wong et al. prepared fluorene-substituted complex Ir159, as well as the homoleptic fac-Ir(C^N)3 complex with the same C^N ligand.174 These complexes have strong absorption in the blue region of the visible spectrum, but no absorption is observed beyond 525 nm. Photoluminescence (PL) for both complexes occurs in the deep red region, 670 nm, with microsecond lifetimes. These complexes were blended with the fullerene acceptor C61-butyric acid methyl ester (PC61BM) and incorporated into bulk heterojunction solar cells. Solar cells doped with Ir159 had modest efficiency, with PCE of 0.14% for two different cells. With the homoleptic Ir(C^N)3 analogue higher open-circuit voltages and short-circuit current densities were obtained, leading to higher PCE values as large as 0.51%, still lower than other cell designs doped with iridium complexes. In work using a similar cell design, Yu et al. doped Ir(piq)2(acac) (Ir151) into ternary bulk heterojunction polymer solar cells, likewise using a fullerene acceptor (PC71BM).175 Complex Ir151 is a close relative of Ir137, Ir148, and Ir149 (Fig. 29), and has nearly identical UV-vis absorption and PL spectra as these other Ir(piq)2(L^X) complexes. Solar cells were fabricated with dopant ratios of Ir151 ranging from 0 to 5 wt%, and especially with low doping percentages the light absorption was dominated by the poly(3-hexylthiophene) (P3HT) polymer component of the solar cell. Significant increases in PCE were observed with iridium doping. An undoped cell had a PCE of 2.99%, which increased to 3.81% with 1 wt% Ir151, and decreased at higher loadings of Ir151. Following a solvent annealing treatment (SAT), an undoped cell had an efficiency of 4.09%, increasing to 4.44% with 1 wt% Ir151. On the basis of PL lifetime measurements, the authors concluded that Ir151 harvested triplet excitons and then efficiently transferred energy to P3HT via a FRET mechanism. The lifetime of a neat film of Ir151 was 3.3 ms, which decreased dramatically to 0.2 ns when Ir151 was doped at 1 wt% into P3HT, indicating 99.9% energy transfer efficiency. It is also possible to directly conjugate Ir(C^N)2(acac) complexes to solar cell polymer materials, and in one study on this strategy dramatic improvements in efficiency were observed. Huang et al. introduced precursor Ir160, which is co-polymerized with two thiophene monomers under Stille conditions to yield doped PTB7 polymer Ir161 (Fig. 33).176 The iridium complex was loaded into the polymer at monomer percentages of 0–5%. The 2-(2,4-difluorophenyl)pyridine (F2ppy) cyclometalating ligand in Ir160 and Ir161 results in very little visible absorption from the iridium complex, and all of the polymers with 0–5% loading of iridium had nearly identical UV-vis absorption spectra. Polymer Ir161 and its undoped analogue were used to fabricate solar cells with a PEDOT: PSS polymer support, Ir161 as the donor, and PC71BM as the acceptor. Two slightly different cell designs were tested, varying the iridium loading and the solar cell cathode materials, and in both cases the Ir-doped cells had higher PCE values, by as much as 45%. All of the cells had very similar open-circuit voltages, and in this case, it was the short-circuit current density that increased upon doping with iridium. The Ir161 polymer emission intensity changed little with and without Ir doping, suggesting that the function of the iridium complex was not to inhibit charge recombination. The precise mechanism for the increase in intensity with Ir loading was not revealed in this study, but this is another example of Ir(C^N)2(acac) complexes positively influencing charge transport in solar cells and leading to improved PCE. To summarize a lot of research described above, Ir(C^N)2(acac) complexes are attractive additives for organic photovoltaic applications. Many of the attributes that make these complexes attractive for OLEDS—their ease of synthesis, high stability, and ability to either solution process or thermally evaporate—likewise are important in solar cell applications. One challenge of the Ir(C^N)2(acac) design is that even with highly conjugated C^N ligands the low-energy charge-transfer absorption bands can only extend to the middle parts of the visible spectrum. Thus, doping these complexes into solar cells is not expected to dramatically improve solar light capture, but as outlined above there are many cases where the long-lived triplet states of these compounds can positively influence charge transport. From a fundamental standpoint continued development of Ir(C^N)2(acac) complexes and related structures for organic photovoltaics and understanding their functions in these devices have merit, but as a widespread technology this class of devices do not hold the same promise as silicon or perovskite-based solar cells.

Organometallic Photosensitizers

313

Fig. 33 Structure of precursor Ir160 and doped conjugated polymer Ir161 (x ¼ 0–0.05).

1.10.5.4.2

Bis-cyclometalated iridium complexes with 2-picolinate ancillary ligands

Another well-known category of charge-neutral bis-cyclometalated iridium complexes feature 2-picolinate ancillary ligands. Some of the earliest work on these compounds varied the conjugation and donor moiety in the 2-substituted pyridine ancillary ligand, showing a change in the nature of the excited state and the color of emission as a function of the ancillary ligand.177 The parent member of this series, Ir(F2ppy)2(2-picolinate) (Ir162, Fig. 34) is often referred to as “FIrpic,” and it has emerged as a very popular sky blue phosphor used in some of the most efficient blue OLED devices.24 Much like the Ir(C^N)2(acac) family, Ir162 has been rarely used as a photosensitizer, though there are scarce reports where it is used in organic photoredox transformations. Yan et al. screened FIrpic (Ir162) as a photocatalyst for a visible-light-induced cascade reaction of isocyanides and N-arylacrylamides with

Fig. 34 Structures of bis-cyclometalated iridium complexes with 2-picolinate ancillary ligands used in photoredox catalysis and DSSCs.

314

Organometallic Photosensitizers

diphenylphosphine oxide.178 They found reasonably good performance with FIrpic in a representative transformation, but ultimately settled on the cationic iridium complex Ir29 as a better alternative. Ir162 was found to be the best photosensitizer for the visible-light mediated nucleophilic addition of a-aminoalkyl radicals to isocyanates or isothiocyanates, outperforming the popular alternatives [Ru(bpy)3]2+, fac-Ir(ppy)3 (Ir4), Ir(ppy)2(acac) (Ir133), and Ir29.179 A number of substrate pairs were screened, partnering substituted N,N0 -dimethylaniline derivatives with aryl isocyanates and aryl isothiocyanates, with good to moderate yields in all cases under compact fluorescent lamp (CFL) illumination. The authors proposed a reductive quenching cycle whereby the FIrpic excited state is reduced by the amine substrate, forming an amine radical cation which deprotonates and adds to the isothiocyanate. The redox-neutral transformation is closed by electron transfer from the FIrpic radical anion to the product radical. Ir162 and other closely related derivatives have also been used on occasion as photosensitizers or charge transport materials in organic solar cells, much in the same manner as Ir(C^N)2(acac) complexes described above. One early report on this application of FIrpic derivatives is from Tian et al., who designed two carboxylate-substituted complexes, Ir163 and Ir164, as dyes for DSSCs.135 These complexes were evaluated alongside cationic complex Ir39, described above. On account of the increased conjugation in the cyclometalating ligands complex Ir164 has the strongest visible absorption of those examined, albeit with minimal absorption beyond 550 nm. Complexes Ir163 and Ir164 have nearly identical photoluminescence maxima, surprising given the difference in C^N ligand, which may indicate that the T1 state is localized on the 5-carboxypicolinate ligand. These complexes were deposited on TiO2 and configured into solar cells using the I−3/I− redox couple. Despite its inferior light absorption, Ir163 performed better as a solar cell dye, with efficiency values about a factor of 1.7 higher than identical cells fabricated with Ir164. However, the efficiency values of these cells were still quite modest, 95% selectivity of CO over H2 produced. Complexes Ir197 and Ir219, a closely related derivative with a different C^N ligand, were investigated along with the analogous bis(imidazoline thione) and bis(imidazoline selenone) complexes, summarized in Fig. 40.208 The UV-vis absorption and photoluminescence spectra have the typical dependence on the C^N ligand, red-shifted in Ir219–Ir221 relative to the C^N ¼ ppy analogues. The thione and selone derivatives have red-shifted and stronger visible absorption than the bis(NHC) complexes, along with photoluminescence that is red-shifted by 560–770 cm−1. These complexes were investigated as photosensitizers for the visible-light-induced oxidative coupling of amines to imines, using O2 as the terminal oxidant. All of Ir197 and Ir217–Ir221 promoted the oxidative coupling of benzylamine, with Ir221 performing the best in this series. The reaction tolerated a variety of substituents at the 4-position of the phenyl ring in the benzylamine substrate and was effective with only 0.25 mol% catalyst loading. The authors demonstrated conversion of 3O2 to 1O2 when Ir220 and Ir221 were irradiated in an aerobic atmosphere, leading them to conclude that an energy transfer pathway, which produces 1O2 that then goes on to induce oxidative coupling, is operative in the catalytic reactions. Notably, complex Ir221, which gave the highest catalytic yields, also produced the most singlet oxygen under identical conditions.

Fig. 40 Bis-cyclometalated iridium complexes with bis(NHC) or bis(imidazoline chalcone) ancillary ligands used for oxidative coupling of benzylamines.

In addition to the symmetric bis(NHC) ligands found in complexes Ir197–Ir221, NHCs have also been paired with pyridine to construct asymmetric neutral donor chelating ligands, which can support bis-cyclometalated iridium photosensitizers. Manzano et al. described the complexes shown in Fig. 41, which include pyridine-NHC ancillary ligands.209 Pyridine-carbene ligands of this type had previously been applied in the design of phosphorescent bis-cyclometalated iridium complexes for light-emitting electrochemical cells,210,211 and in this work these complexes were applied as photosensitizers for photocatalytic water reduction

Fig. 41 Bis-cyclometalated iridium complexes with pyridine-NHC ancillary ligands.

Organometallic Photosensitizers

321

to hydrogen. The collection of complexes in Fig. 41 vary with respect to the cyclometalating ligand, either ppy or F2ppy, as well as the substitution pattern on the pyridine-NHC ligand. The photoluminescence spectra of the F2ppy complexes (Ir229–Ir233) were quite similar to one another, in the blue region each with two vibronically structured maxima at 451(1) and 477(1) nm, whereas in the ppy complexes (Ir222–Ir228) there were two different spectral profiles. In complexes where the pyridine-carbene ligand lacks a methyl spacer and has methyl substituents on the pyridine (Ir222, Ir224, and Ir225) the photoluminescence has vibronic structure and occurs in the sky-blue to blue-green region, lmax ¼ 472(1) and 501(1) nm. In the remaining ppy complexes, where the pyridine ring has a nitro substituent or a methylene spacer between the pyridine and carbene, the photoluminescence is more reminiscent of a pure charge-transfer excited state, with a broad featureless peak in the green region, centered at 512(2) nm. In terms of the redox properties, the substituent pattern on the pyridine-NHC ancillary ligand had only subtle effects, whereas in general the C^N ¼ F2ppy complexes were 0.2–0.3 V more difficult to oxidize and had HOMO–LUMO gaps 0.2–0.3 eV larger than their ppy analogues. The T1 excited state in these compounds was determined by DFT to involve 3MLCT character, mixed with either 3 LC or 3LLCT character. Some of these complexes along with closely related reference compounds210,211 were examined as photosensitizers for water reduction, using [Co(bpy)3]2+ as the catalyst and triethanolamine as the sacrificial reductant, in acidified water/MeCN solvent. During 2 h of irradiation, complex Ir223 gave the largest amount of hydrogen, 13 TON relative to the photosensitizer. On the basis of Stern-Volmer quenching experiments, the authors concluded that an oxidative quenching pathway was likely operative, with the excited photosensitizer donating an electron to the cobalt catalyst. The abovementioned complexes all used the ubiquitous imidazoline NHC as part of their ancillary ligand, but it is also possible to prepare cationic bis-cyclometalated iridium complexes with less typical N-heterocyclic carbene ancillary ligands. Lam et al. described the series of complexes summarized in Fig. 42, where the pyridinium-derived NHC is installed via reaction with N-(2pyridyl)-pyridinium salts.212 Related complexes and their photoluminescence properties had been previously studied, where the pyridinium-derived chelate served as the main cyclometalating ligand, and in the compounds in Fig. 41 the pyridinium-derived ligand is the ancillary ligand, partnered with the common cyclometalating ligands 2-phenylpyridine and 1-phenylpyrazole. The substituents on the ancillary ligand were varied to include nitrogen-donor substituents, which can donate into the pyridinium via resonance delocalization, as well as tert-butyl or aryl substituents where that is not possible. The emission spectra are highly sensitive to the substituent pattern on the pyridylidene and much less so to the identity of the cyclometalating ligand, suggesting that the luminescence arises from a 3MLCT state involving the ancillary ligand. Consistent with this notion, cyclic voltammetry indicates that the LUMOs in dimethylamino-substituted complexes Ir238 and Ir243 are highly destabilized and their photoluminescence profoundly blue-shifted, relative to tert-butyl-substituted Ir234 and Ir242. Complex Ir234 is a strong photooxidant, with the excited-state potential E( IrIII/IrII) ¼ +0.71 V vs. Fc+/Fc, significantly more oxidizing than typical cyclometalated iridium complexes and even [Ru(bpy)3]2+. The oxidizing nature of this complex’s excited state was demonstrated in photocatalytic radical thiol-ene additions, which involves oxidation of the thiol substrate. In contrast, dimethylamino-substituted Ir238 is a reasonably strong photoreductant, E(IrIV/ IrIII) ¼ −1.87 V vs. Fc+/Fc, allowing it to be used as a photosensitizer for visible-light-driven CO2 reduction to CO, in concert with a pentadentate cobalt catalyst. Under blue LED irradiation with triethylamine as the sacrificial reductant, 1900 TON of CO were produced in 75 h, higher than analogous systems which use fac-Ir(ppy)3 as the photoreductant.

Fig. 42 Bis-cyclometalated iridium complexes with pyridylidene ancillary ligands.

As mentioned above, complexes with NHC-based C^C: cyclometalating ligands are rarely suitable as visible-light photosensitizers, since the large HOMO–LUMO gaps result in absorption primarily or exclusively in the UV region. However, a recent report by Sun et al. shows that it is possible to incorporate BODIPY moieties onto a cyclometalated NHC, leading to complexes with strong visible absorption from the BODIPY.213 These heteroleptic complexes are displayed in Fig. 43, and they all include two benzo[h] quinoline cyclometalating ligands along with the BODIPY-substituted NHC. The photophysical properties of these complexes represent a sum of their parts. The UV-vis absorption is dominated by the BODIPY, with an intense band that occurs between 530 and 543 nm, with high extinction coefficients 8  104 M−1 cm−1 in each case. The photoluminescence in these compounds is a mirror image of the absorption, has a short ns-scale lifetime component, and is minimally quenched by O2, leading to its assignment as BODIPY fluorescence. However, the authors also observed PL decay on the ms timescale, suggesting that the normal

322

Organometallic Photosensitizers

Fig. 43 Charge-neutral heteroleptic cyclometalated iridium complexes with BODIPY-substituted NHC ligands. 3

CT phosphorescence from the iridium complex overlaps with this fluorescence. The photoluminescence occurs at longer wavelength, 610 nm, in complexes Ir244 and Ir245, where the BODIPY is substituted onto the NHC phenyl ring with an alkyne spacer; the remaining compounds have their PL peak at 583–587 nm. Complexes Ir245 and Ir248, with water-solubilizing oligoether substituents, were examined as photosensitizers for singlet oxygen generation. In these studies there was some evidence of synergism between the components. BODIPY chromophores on their own do not produce singlet oxygen, since intersystem crossing into the triplet manifold does not occur. However, both of complexes Ir245 and Ir248 produce modest amounts of singlet oxygen when excited in the near-UV or visible region. In complex Ir245 the quantum yield (FD) was 0.38 with 450-nm excitation and 0.37 at 534 nm, whereas in Ir248 a lower quantum yield was observed, maximizing at 0.22 with 450-nm excitation. Both complexes had increased toxicity towards cancer cells upon light activation, suggesting their potential as photodynamic therapy agents, although the toxicity did not correlate with FD. This led to the conclusion that other reactive oxygen species (ROS) were responsible for the photocytotoxicity.

1.10.5.6

Bis-cyclometalated iridium complexes for enantioselective photoredox transformations

All of the tris- and bis-cyclometalated iridium complexes described above belong to chiral rotational point groups, but in the vast majority of applications they are isolated and used as racemic mixtures. Most enantioselective photoredox methodologies use a racemic photosensitizer in combination with a chiral co-catalyst that directs the enantioselectivity.214 However, there has been some recent success in isolating enantiopure bis-cyclometalated iridium complexes, in some cases supported by chiral ligands, to promote enantioselective photochemical transformations at a single site. The first success in this area came from Meggers et al., who reported two closely related chiral-at-metal complexes, Ir249 and Ir250.215 These complexes, shown in Fig. 44, were originally prepared as

Fig. 44 Chiral-at-metal bis-cyclometalated iridium complexes for enantioselective photoredox transformations.

Organometallic Photosensitizers

323

chiral Lewis acid catalysts, and they can be accessed in two steps from the chloro-bridged dimers, [Ir(C^N)2(m-Cl)]2.216 The dimeric starting materials are isolated in racemic form, i.e. a 50:50 ratio of D and L forms, and treatment with a chiral auxiliary ligand generates a diastereomeric mixture of products, separable by column chromatography. From there, protonolysis of the diasteromerically pure L intermediate in the presence of MeCN produces the enantiopure complexes. In the initial report on the photoredox applications of these complexes, the first reaction described was the visible-light-promoted enantioselective alkylation of an acyl imidazole, using substituted benzyl bromides as the reaction partner. With both enantiopure catalysts high product yields and e.e. >95% were observed. The complex Ir250 was found to be the faster of the two catalysts investigated, and with this catalyst a wide range of acyl imidazole and benzyl bromide substrates were tolerated, provided the benzyl bromide was substituted with one or more electron-withdrawing groups at the 2- and/or 4-positions. The proposed mechanism includes coordination of the acyl imidazole in its deprotonated enolate form, with concomitant photoredox generation of a benzyl radical which adds stereospecifically to the coordinated enolate. Thus, the iridium catalyst serves the dual roles of coordinating one substrate and directing stereospecific radical addition, while also serving as a photoreductant to generate the radical partner. The Meggers group has continued to develop other asymmetric photoredox transformations using Ir249, Ir250, and the closely related complex Ir251 which has an additional tert-butyl substituent on the C^N ligand. These enantioselective transformations include trichloromethylation217 or perfluoroalkylation218 of 2-acyl imidazoles, and cross-coupling reactions for diastereoselective synthesis of 1,2-amino alcohols,219 which proceed by similar mechanisms involving direct substrate coordination to the iridium center. With a slightly different strategy, this group has also developed a sequential reaction involving asymmetric hydrogenation followed by photoredox radical trihalomethylation/cyclization.220 In this approach the ketone hydrogenation substrate is directed to the chiral metal center by a secondary ligand 2,4-dimethylpyrazole (dmp), included as an additive. Hydrogenation proceeds thermally, and then photoredox chemistry generates the trichloromethyl or trifluoromethyl radical that initiates the radical addition/cyclization cascade. Complexes Ir249–Ir251 described above include bound MeCN ligands which can be readily displaced, allowing the substrate(s) to bind at the chiral metal center. Another complementary strategy for asymmetric photoredox catalysis is to prepare enantiopure catalysts which include a functional chelating ligand that is not displaced during catalysis but directs the substrate in some way for an asymmetric transformation. Yoon et al. have designed chiral complexes Ir252–Ir254, with differing degrees of fluorination on the cyclometalating ligands (Fig. 45).221 The role of the pyridylpyrazole ancillary ligand is to position the substrate via hydrogen bonding for the enantioselective photoreaction. They showed that these compounds promote enantioselective [2 + 2] photocycloaddition reactions of 3-alkoxyquinolone substrates, where the alkoxy group includes a tethered alkene that participates in the cycloaddition. These reactions proceed via triplet Dexter energy transfer from the iridium complex to the substrate, which is bound to the catalyst via a combination of hydrogen bonding and p–p interactions. Complex Ir254 bound the substrate the strongest, and thus gave the highest enantioselectivity. In subsequent work on a similar catalyst system, Yoon et al. also described variants Ir255–Ir257, varying the substitution on the C^N and/or N^N ligands.222 In this work they explored enantioselective

Fig. 45 Bis-cyclometalated iridium complexes with pyridylpyrazole ancillary ligands, used for [2 + 2] photocycloaddition reactions.

324

Organometallic Photosensitizers

intermolecular [2 + 2] cycloaddition reactions between quinolone substrates and maleimide. Catalyst screening in this system showed that variant Ir257 gave the highest yields and highest enantioselectivities, and this catalyst was applied with a number of different quinolone substrates, as well as various N-substituted maleimides. Excited-state quenching studies showed that even though the catalyst’s pyrazole moiety binds tightly with the quinolone through a bidentate hydrogen bonding interaction, energy transfer to the quinolone is slow. Instead, they found that the catalyst-quinolone pair is preferentially quenched by the maleimide, and that the triplet-state maleimide undergoes a stereospecific addition to the bound quinolone. Thus, in this transformation it is not necessary that the triplet energy acceptor, the maleimide, actually be bound tightly by the chiral catalyst. In addition to ligands that direct substrates to the chiral catalyst via hydrogen bonding, there have also been strategies developed where the chiral iridium catalyst is supported by a chiral ligand, which directs stereospecificity of the photochemical transformation. One approach along these lines was introduced by Baslé et al., who prepared all four diastereomers of a bis-cyclometalated iridium complex supported by a chiral NHC-carboxylate ancillary ligand.223 These structures are summarized in Fig. 46, and they were accessed by treating racemic [Ir(ppy)2(m-Cl)]2 (Ir1) with a slight excess of the enantiopure NHC-carboxylate. With using the (S) enantiomer of the ancillary ligand L-(S)-Ir258 was formed preferentially in a 2:1 ratio over the D-(S) diastereomer, and with the (R) enantiomer of the ancillary ligand the D-(R)-Ir258 product is the major species formed. This diastereomeric selectivity also allowed kinetic resolution of the chloro-bridged dimer Ir1. By using 0.5 equivalent of the NHC-carboxylate, only one enantiomer of Ir1 reacted to form the preferred diastereomer of Ir258, leaving behind the other enantiomer of Ir1, which could be isolated and, as a proof of concept, converted to enantiopure Ir(ppy)2(acac) (Ir133). Using either L-(S)-Ir258 or D-(S)-Ir258 as the photosensitizer, stereospecific [2 + 2] photocycloadditions of cinnamic acid and 1,1-diphenylethylene were conducted, producing a substituted cyclobutene product with >98:2 diastereoselectivity in racemic form. The reaction was proposed to involve energetically favorable triplet energy transfer from 258 (ET1 ¼ 58.8 kcal/mol) to cinnamic acid (ET1 ¼ 57.4 kcal/mol).

Fig. 46 Chiral bis-cyclometalated iridium complexes with NHC-carboxylate ancillary ligands.

Ye et al. have applied enantiopure bis-cyclometalated iridium complexes in a photochemical strategy for kinetic resolution of racemic amino acids.224 They started with complex D-Ir259 (Fig. 47), an analogue of Ir249 with C^N ¼ pq and D stereochemistry, which reacted with racemic amino acids to form complexes Ir260–Ir266. These amino acid complexes can be isolated,225 but the bound amino acids in the L configuration are readily photooxidized to the corresponding imino acids in the presence of oxygen, allowing imino acid complexes Ir267–Ir273 to be prepared. Complexes where the amino acid is in the D configuration do not oxidize under the same conditions, enabling a resolution strategy whereby racemic amino acids are combined with D-Ir259 in the presence of base and O2, and during photolysis a mixture of the imino acid complex and the D-amino acid complex are formed; these two complexes can be readily separated chromatographically, and the D-amino acid can be cleaved from the iridium center using HCl. In all cases the deracemization reactions gave yields >90% and diastereomeric excess of >99% for the D-D-Ir260–Ir266 complex, in 1–3 cycles. The photooxidation involves production of singlet oxygen, which then is the active species for dehydrogenation of the coordination amino acid.

Organometallic Photosensitizers

325

Fig. 47 Bis-cyclometalated iridium amino acid and imino acid complexes.

The ability to promote enantioselective photoredox transformations with cyclometalated iridium complexes is a major development for synthetic methodology. Most photoredox methodology either is not stereospecific, proceeding through SET processes that offer no control over stereochemical outcomes, or for certain types of transformations requires stereoselective co-catalysts to promote enantioselectivity. However, the development of chiral photosensitizers which on their own can give rise to enantioselection comes with a unique set of challenges, and it is unclear at this early stage how widespread in applicability this approach will be. In general outer-sphere photoinduced SET does not require a specific orientation of the photosensitizer and quencher to proceed. As a result, photosensitizers that are involved in stereospecific photoredox steps require tight binding of the substrate, either directly to the chiral metal center or through noncovalent association with a chiral supporting ligand, which can be seen from the examples highlighted above. Thus, the substrates used for these reactions must have properly oriented functional groups that permit the interactions with the photosensitizer that result in stereospecific electron or energy transfer. There remains a need for additional approaches to this problem, where a photosensitizer serves the dual function of mediating excited-state chemistry while also guiding the stereochemical outcome.

1.10.5.7

Bis-cyclometalated iridium complexes with monodentate ancillary ligands

Although there are a few select examples given above of bis-cyclometalated iridium complexes with monodentate ancillary ligands, the vast majority of cyclometalated iridium photosensitizers are tris-chelated structures, with three bidentate ligands. However, there are two classes of complexes with monodentate ancillary ligands which have received considerable attention: anionic complexes of the general formula [Ir(C^N)2(X)2]−,226 where X is an anionic ligand, usually CN− or SCN−, and cationic complexes of the type [Ir(C^N)2(CNR)2]+, where CNR is an isocyanide.227–233 These previous reports primarily investigated the phosphorescence of these compounds, and in some cases applied them in optoelectronic devices. However, there are a few reports that investigate complexes in these classes as photosensitizers. As part of their study on photoinitiators for polymerization, Lalevée et al. screened anionic bis-cyano complexes Ir274–Ir276 (Fig. 48), none of which were active for light-initiated cationic polymerization of an epoxide, but complex Ir274 did show modest activity in a radical acrylate polymerization.167 In a study on singlet oxygen generation, Zhao et al. studied the tetrabutylammonium salt of Ir276, the bis-terpyridine complex Ir277, as well as the double salt Ir278 which pairs the cation and anion. The blue photoluminescence of Ir276 overlaps with the MLCT absorption of the cation Ir277, and Stern-Volmer quenching studies revealed that the phosphorescence from Ir276 was quenched when the cation was added. Irradiating with white light in the presence of oxygen, the authors found that the double salt Ir278 produced singlet oxygen 4.2 times faster than Ir277 on its own, and 2.5 times faster than the anion Ir276 on its own. The authors suggested that energy transfer between the anion and cation is responsible for this increase in 1O2 generation. In a separate experiment, they studied singlet oxygen generation with varying ratios of Ir276 and Ir277 and found an increase in rate up to the charge-balanced 1:3 ratio of the two, with no further increase beyond 3 equivalents of the cation.

326

Organometallic Photosensitizers

Fig. 48 Anionic bis-cyclometalated iridium complexes used as photosensitizers.

There is one report that investigated singlet oxygen photosensitization by a cationic [Ir(C^N)2(CNR)2]+ complex.234 Complex Ir279 (Fig. 49), along with its chloro-bridged dimer and several rhenium, ruthenium, and osmium diimine complexes, were investigated as photosensitizers for singlet oxygen generation. The phosphorescence of Ir279 is centered at 508 nm with a lifetime of 2.70 ms, and it is quenched in the presence of O2. The authors measured a singlet oxygen quantum yield, FD, of 0.54 for Ir279, over twice as high as the chloro-bridged dimer. In their study of Ir279 they found that 95% of the triplet excited states were quenched by oxygen in air-equilibrated MeCN solution, with 57% of those quenching events leading to the production of 1O2.

Fig. 49 Structure of Ir279, a cationic bis-cyclometalated iridium bis-isocyanide complex used as a singlet oxygen photosensitizer.

Organometallic Photosensitizers

1.10.5.8

327

Cyclometalated iridium complexes with tridentate ligands

All of the structures presented above in this section on organometallic iridium photosensitizers feature two or three bidentate cyclometalating ligands. In an effort to further improve the thermal and photostability of cyclometalated iridium photosensitizers, and/or diversify their photophysical properties, several recent efforts have focused on analogues where the cyclometalating and/or ancillary ligands are tridentate, meridionally-coordinating ligands. Some of the earliest work on tridentate cyclometalated iridium complexes focused on bis(tridentate) structures, where the tridentate ligands are of the C^N^C, N^C^N, C^N^N, and N^N^N varieties, arranged in some combination around the iridium(III) centers. These early works focused on the photophysics and redox properties,235–241 and in one case on the assembly of these complexes into linear bimetallic arrays with coupling chemistry.238 In one of the first reports to investigate tridentate cyclometalated iridium complexes as photosensitizers, Tinker and Bernhard prepared the homoleptic complex [Ir(phbpy)2]+ (Ir280, phbpy ¼ 6-phenyl-2,20 -bipyridine). The structure of this compound is shown in Fig. 50, and it has an identical coordination environment as the ubiquitous photosensitizer [Ir(ppy)2(bpy)]+ (Ir23), except with two tridentate ligands instead of three bidentate ligands. The visible absorption of Ir280 is about twice that of Ir23, as measured by the molar absorptivity at 460 nm, whereas the photoluminescence wavelength, quantum yield, and lifetime are similar. In addition, the redox potentials of these two complexes are within 20 mV of each other. This complex was used as a photosensitizer for visiblelight-driven hydrogen evolution, in different organic solvent/water mixtures with K2PdCl4 as the hydrogen evolution precatalyst and triethylamine as the sacrificial reductant. In MeCN/water, the initial rate of hydrogen evolution was slightly smaller with Ir280 compared to Ir23, but catalysis was much more robust with the bis-tridentate complex Ir280, leading to turnover numbers almost three times as large. The results were similar but not as dramatic in DMF/water, and in THF/water Ir23 evolved almost twice as much H2, and it was in this solvent system that catalysis was most robust.

Fig. 50 Cyclometalated iridium complexes with tridentate ligands used as photosensitizers for hydrogen evolution, CO2 reduction, and photoredox catalysis.

328

Organometallic Photosensitizers

Another class of tridentate cyclometalated iridium complexes which have become prominent in photocatalytic applications have the general structure [Ir(C^N)(N^N^N)X]+, where N^N^N is a terpyridine derivative, and X is a monodentate anionic ligand. These were first introduced by Bernhard et al., and in addition to a detailed study of their photophysical and electrochemical properties, these photosensitizers were used for photocatalytic hydrogen evolution, in MeCN/water with colloidal platinum as the catalyst and triethylamine as the sacrificial reductant.242 In this work they studied chloride-terminated complexes Ir281–Ir287, as well as two cyano analogues Ir288 and Ir289. In the chloride series the formal IrIV/IrIII redox couple depends primarily on the structure of the C^N ligand. The oxidation potentials of C^N ¼ ppy complexes Ir281 and Ir284 are nearly identical, shifting anodically by 60 mV in the fluorinated derivatives Ir282 and Ir285, with methoxy substitution inducing a 260–270 mV cathodic shift in Ir283 and Ir286. In analogue Ir287 the first oxidation occurs on the electron-rich substituted terpyridine ligand, and in the cyano complexes Ir288 and Ir289 the oxidation potential is 180–260 mV more positive at parity of C^N and N^N^N ligands. The first reduction potentials of these complexes span a very narrow range of −1.45 to −1.50 V vs. Fc+/Fc, indicating a terpy-centered LUMO whose energy depends little on the substitution pattern. The chloride-substituted complexes luminesce in the green to yellow region, with cyano-substitution inducing a 45 nm (1800–2000 cm−1) blue shift. These complexes are all weaker photoreductants than reference compound Ir23, by 120–330 mV, but they are all stronger photooxidants by at least 450 mV. In MeCN/water mixtures these photosensitizers typically outperform Ir23 as photosensitizers for hydrogen evolution, with the cyano complexes Ir288 and Ir289 appearing to be the most robust and giving the highest TONs. This observation suggests that the tridentate nature of these complexes prevents ligand substitution decomposition pathways that commonly occur in strongly coordinating solvents like MeCN. To further optimize the photophysical properties within this structural class, a series of “push-pull” derivatives were synthesized, which partner an electron-rich tert-butyl-substituted terpyridine with a fluorinated 2-phenylpyridine derivative (Ir290–Ir298, Fig. 50).243 These complexes were prepared from the Ir(N^N^N)Cl3 precursor by either CdH or CdF activation of the 2-phenylpyridine proligand derivative. The spatially separated frontier orbitals in these compounds are a HOMO localized on the Ir d orbitals and the aryl ring of the C^N ligand, and a LUMO localized on the terpyridine with some Ir character. The charge-transfer absorption bands and phosphorescence blue shifts with increasing fluorination of the C^N ligand, and replacing the chloro ligand with a cyano likewise induces a blue shift. The IrIV/IrIII redox potentials span a range of 1.21–1.60 V vs Fc+/Fc, shifting more positively as fluorination is increased and when X ¼ CN. The first reduction potentials of all complexes span a much narrower range, consistent with the observation that the LUMO is primarily terpy-centered. In terms of their excited-state redox potentials, the reducing and oxidizing power of these compounds fall in a similar range as those of Ir281–Ir289, making them weaker photoreductants than Ir23, but stronger photooxidants. These complexes were evaluated as photosensitizers for photocatalytic hydrogen evolution, in 8:1 MeCN/water with in-situ-generated colloidal Pt as the catalyst, and triethylamine as the sacrificial reductant. Complex Ir297 performed the best and represented a slight improvement in turnover number over previously described Ir287.242 Complexes Ir291 and Ir296 were also evaluated as photosensitizers for an organic photoredox transformation, the decarboxylative fluorination of carboxylic acids, using Selectfluor® as the fluorine source. Both of these complexes dramatically outperformed the well-known organic photoredox catalyst Ir32. Complexes in this family are also known to be photocatalysts for CO2 reduction, operating as standalone catalysts where light absorption and chemical reduction events both occur at the iridium complex. This was first demonstrated by Sato et al., using the unsubstituted complex Ir281.244 Reduction of CO2 is proposed to involve a combination of outer-sphere electron-transfer quenching, where the excited iridium complex reacts with triethanolamine to form the one-electron reduced species, and inner-sphere reactions that involve replacement of the Cl– ligand with a hydride and subsequent formation of ill-defined CO2 adducts. In this initial report the reaction was carried out in CO2-saturated MeCN/TEOA (5:1), and the reaction was >98% selective for CO over H2 and HCOOH. A TON of 38 for CO generation was observed, and the quantum yield for CO generation was 0.13 when irradiated at 480 nm. The authors also investigated two substituted analogues of Ir281 with either CH3 or CF3 substitution on the ppy ligand (the location of the substituent was not specified), and found slightly better performance with the more electron-rich methyl-substituted analogue, but less CO production with the CF3 analogue. From this, they concluded that the hydricity of the intermediate iridium hydride was an important factor for catalysis. To improve the performance of these complexes, Bernhard et al. placed aryl substituents on the equatorial position of the terpy ligand, in a set of complexes that includes Ir284 and new complexes Ir299 and Ir300.245 The aryl substituents have a trivial influence on the reduction potentials, and the absorption wavelengths and molar absorptivities are all similar as well. The effect of the anthryl substituent in Ir299 is to localize the lowest triplet state on the anthryl moiety, which has the effect of lowering the photoluminescence quantum yield via energy transfer to the nonradiative anthryl state. These complexes were tested as photocatalysts for CO2 reduction to CO, using white light or blue LED irradiation, under otherwise identical conditions as the original report where unsubstituted analogs were used.244 They found an improvement in the TON of CO when using lower intensity blue light, which inhibits catalyst deactivation pathways, and by slowing down the reaction they could achieve larger overall TON. Complex Ir299 had the smallest initial TOF for CO production, again showing that electron-withdrawing groups slow down catalysis, but the anthryl-substituted analogue was proposed to provide greater steric protection for longer-lasting catalysis and higher TON. Using the same structural class, Gasser et al. designed complex Ir301, which is capable of localizing in the mitochondria and doing intracellular photoredox catalysis.246 The 3-phenylisoquinoline cyclometalating ligand in Ir301 extends the absorption beyond 500 nm, and produces stable photoluminescence in the yellow region. Aerobic photolysis of Ir301 in the presence of NADH gave 95% yield of NAD+ after 30 min, showing that complex Ir301 can photooxidize this important redox-active biomolecule. In the absence of oxygen, NAD• radicals were formed in the presence of the biological oxidant cytochrome c oxidase, which serves as the terminal oxidant in anaerobic conditions. This photooxidation of NADH both in the presence and absence of oxygen results in

Organometallic Photosensitizers

329

phototoxicity to both normoxic and hypoxic cells, localizing in the mitochondria and disrupting electron transport pathways. This strategy allows for oxygen-independent photodynamic therapy, a significant challenge in cancer therapeutics. Complexes of the [Ir(C^N^C)(N^N^N)]+ structure type have also been tested as dyes for dye-sensitized solar cells.247 The carboxylate-terminated photosensitizers shown in Fig. 51 were all prepared and used in nanocrystalline TiO2-based dye-sensitized solar cells. These complexes all have strong absorption in the blue region of the visible spectrum, arising from charge-transfer bands which are similar in wavelength but most intense for complex Ir303, which has pendant aryl substituents on both the N^C^N and N^N^N ligands. These compounds are photoluminescent in the deep red region, and in complexes Ir302 (l ¼ 679 nm) and Ir303 (l ¼ 685 nm) the emission is at a longer wavelength than in Ir304 (l ¼ 654 nm) and Ir305 (l ¼ 653 nm), on account of the extended conjugation of the N^N^N ligand in the former. These complexes were compared to Ir37, which had previously been used in DSSCs but suffered from poor visible light absorption.132 Compared to this reference compound, Ir302–Ir305 all have much stronger visible absorption, with similar HOMO energies and LUMOs that are stabilized by 0.3–0.4 eV. Complex Ir305 had the highest solar cell efficiency of this series, with a PCE of 2.16%, almost two times higher than the rest on account of both a larger short-circuit current density and a higher open-circuit voltage. However, this complex still fell well short of the well-known N3 dye used in otherwise identical cells, which gave a PCE of 3.32% on account of a much larger short-circuit current density.

Fig. 51 Carboxylate-terminated [Ir(C^N^C)(N^N^N)]+ complexes used as solar cell dyes.

1.10.5.9

Summary and outlook

It is clear from the above discussion in this section that iridium(III) complexes represent the most widely studied and successful class of organometallic photosensitizers, impacting many applications where photosensitizers are important. There is a huge structural variety already known, and still opportunities to explore other ligand designs in the discovery of next-generation analogues. Many of these compounds had already found great success in OLEDs, and their photostability, long-lived charge-transfer excited states, and versatile photoinduced redox chemistry have likewise made them a mainstay in photoredox catalysis. The long-known homoleptic fac-Ir(ppy)3 complex has been particularly prominent in challenging reductive transformations, although synthetic challenges with this class of compound have limited its development. More recently many heteroleptic designs have emerged that provide even greater control over the excited-state redox chemistry. In particular, the [Ir(C^N)2(N^N)]+ family is a very versatile platform that can carry out less demanding reductive transformations but can also be tuned to be effective for oxidative elementary steps, and charge-neutral Ir(C^N)2(L^X) complexes, where L^X is an anionic chelate, have begun to become prominent for challenging reductions. Once major inherent limitation of iridium is its cost, being one of the least abundant and most expensive transition metals. The high cost of iridium is not a particular detriment to organic photoredox methodology, where the products are typically high-value molecules and the most important consideration is whether those molecules can be made efficiently and in high yield. However, for applications like solar fuels and photovoltaics, which would need to be run on huge scales to be truly impactful to society, the high cost of iridium is problematic. Another issue with most iridium photosensitizers that has plagued their development in solar energy applications is their poor absorption over most of the solar spectrum, typically anything beyond the blue region. Several groups have tackled the problem of increasing visible light absorption, with many of those examples highlighted above. However, changes to the cyclometalating ligand typically result in only incremental advances in light absorption, and where groups have been successful in producing panchromatic absorption by judicious ancillary ligand design, e.g. Ir84–Ir94 in Fig. 20, the lifetimes are too short and/or the excited-state redox potentials are poorly matched to typical semiconductor substrates. As a result, it still remains an outstanding fundamental challenge to design cyclometalated iridium complexes that both absorb over most of the visible spectrum and have lifetimes and excited-state potentials suitable for charge transport. A successful design that meets both of these criteria will likely still have limited impact in solar cell design, given the high cost of iridium. That said, the vast majority of photosensitizers

330

Organometallic Photosensitizers

used in photoredox catalysis only absorb in the blue region, so analogues with extended absorption may have value. Complexes that tolerate longer excitation wavelengths may be useful in more complex photoredox schemes where wavelength selection can be used to drive two or more different photosensitizers for orthogonal transformations in the same reaction pot.

1.10.6

Rhodium photosensitizers

Using the 5-(tert-butyl)-2-phenyl substituted benzoxazole and benzothiazole ligands, Meggers prepared two chiral-at-metal rhodium complexes (Rh1 and Rh2) that are analogous to Ir249–Ir251 and studied their photocatalytic performance (see Fig. 52). These complexes were also synthesized from the chloro-bridged Rh dimers with a chiral auxiliary ligand, followed by treatment with NH4PF6 in acetonitrile, and they contain two labile acetonitrile ligands that can be readily replaced by organic substrates in asymmetric photoredox catalysis.248 Similar to Ir249–Ir251, these Rh complexes can function as both Lewis acid catalysts and as photocatalysts to promote a series of enantioselective visible-light-mediated photoredox catalysis reactions.217,248–252 As a continuation of this family of compounds, the same group reported several other chiral-at-rhodium complexes with Rh3 and Rh4 being representative examples, summarized in Fig. 52. To investigate the effect of cyclometalating ligands on the asymmetric catalysis, Rh3 was prepared with a tert-butyl substituted phenylindazole bis-cyclometalating ligand.253 This complex was tested as a dual functional catalyst for a-cyanoalkylation of 2-acyl imidazoles. When comparing to its analogues Ir250 and Rh2, Rh3 demonstrates significantly higher reaction yield with comparable enantioselectivity with respect to Ir250 due to the slower ligand exchange kinetics at the Ir center, and it is superior to Rh2 in terms of both yield and ee. Complex Rh4 was designed with two different cyclometalating ligands, which are a benzimidazole and a benzothiazole, also partnered with two acetonitrile ligands.254 This complex was tested for asymmetric [2 + 2] photocycloadditions with a,b-unsaturated N-acyl pyrazoles and alkenes using blue LED irradiation. More importantly, this provides a promising approach for synthesizing related chiral-at-metal complexes with two different cyclometalating ligands to further modify their photophysical and photocatalytic activity.

Fig. 52 Organometallic rhodium photosensitizers.

Another notable example of organometallic Rh photosensitizer is a biphenyl complex Rh5 included in Fig. 52. This complex is identified as a triplet sensitizer with triplet excited energy of 53 kcal/mol and shows phosphorescence at lmax ¼ 540 nm with a quantum yield of 0.14 at room temperature.255,256 This complex’s excited-state lifetime of 181 ms in toluene and 338 ms in 2-MeTHF is significantly longer than the ubiquitous fac-Ir(ppy)3 and Ru(bpy)2+ 3 . These photophysical attributes made it an excellent photosensitizer for CdF borylation of fluoroarenes with B2pin2 in the presence of Ni(IMes)2 (IMes ¼ 1,3-dimesitylimidazoline2-ylidene) as the co-catalyst. The catalytic transformations proceed through energy transfer from Rh5 to the Ni catalyst after undergoing oxidative addition with the fluoroarenes. Although Rh5 possesses excited-state redox potentials similar to those of fac-Ir(ppy)3 [E(RhIV/ RhIII ¼ −1.73 V vs. Fc+/Fc), E(RhIII/ RhII ¼ 0.04 V vs. Fc+/Fc)], SET in the catalytic cycle was excluded.

Organometallic Photosensitizers

1.10.7

331

Palladium photosensitizers

Palladium complexes are well known as versatile catalysts in organic synthesis such as Suzuki and Heck coupling reactions. Photosensitizers based on organometallic palladium complexes are sparse mainly due to the small ligand field splitting of palladium rendering the metal-centered d-d states thermally accessible, leading to poor luminescence and short-lived excited states at room temperature. Thus, they typically contain strong donor ligands to raise the d-d transition energies. Ong et al. reported Pd1, which contains a strong s-donating carbodicarbene pincer ligand, and explored its application in photoinduced catalysis (see Fig. 53).257 This complex exhibits an intense absorption at l ¼ 499 nm and luminesces at l ¼ 570 nm, with excitations from both LC and MLCT charge-separated states, and is a versatile catalyst for both thermal and photochemical transformation. Complex Pd1 promoted photocatalytic CdN coupling of a series of aldehydes and amines under visible light in the presence of oxygen. In this reaction photoreduction of O2 to H2O2 was proposed, with the hydrogen peroxide then participating in an oxidative CdN coupling between the aldehyde and the amine. Another photocatalytic reaction promoted by Pd1 is CdH arylation, which combines a photoredox cycle and a more typical cross coupling cycle. Photoredox catalysis generates aryl radicals from diazonium precusors, which then couple with the 2-phenylpyridine derivative coupling partner, via coordination to Pd followed by reductive elimination. In addition, this complex in the ground state catalyzes the Suzuki-Miyaura cross-coupling reaction of aryl bromides and phenyl boronic esters. Combining the photocatalytic CdN coupling with the ground-state Suzuki-Miyaura coupling, Pd1 can be used as a dual functional catalyst for photocatalytic C–H arylation in a one-pot setup.

Fig. 53 Organometallic palladium photosensitizers.

Che et al. introduced a suite of Pd complexes (Pd2–Pd9) supported with tetradentate O^N^C^N ligands, also shown in Fig. 53.258 These ligands are constructed by having a phenolate tethered to one of the pyridine rings in the conventional N^C^N cyclometalating ligands. This series of complexes show phosphorescence in the green region with lmax ranging from 498 to 543 nm in dichloromethane solutions at room temperature. Complexes Pd5–Pd9 display significantly improved quantum yields and prolonged excited-state lifetimes with respect to Pd2–Pd4 that can be assigned to the suppression of ligand deformation after photoexcitation, evidenced by DFT calculations. Complex Pd8 has the highest quantum yield of 0.22 while Pd7 shows the longest excited-state lifetime of 122 ms among these complexes. In the same study the authors investigated the photoinduced energy transfer of Pd5–Pd9 to 9,10-diphenylanthracene (DPA) and observed the generation of 3 DAP, followed by triplet-triplet annihilation giving rise to the singlet excited species (1 DAP). Complexes Pd5–Pd9 were also applied as photosensitizers to facilitate oxidative CdH functionalization of secondary amines with O2, which was initiated after the formation of singlet oxygen as the active oxidant by the photosensitizers under light excitation.

332

1.10.8

Organometallic Photosensitizers

Platinum photosensitizers

Even though the majority of metal-based photosensitizers are Ir and Ru complexes, organoplatinum complexes in photosensitization have also received considerable attention. The unique square planar geometry of Pt complexes with vacant axial coordination sites could potentially allow for both outer-sphere and inner-sphere interactions with substrates, rendering them promising to be used in catalytic transformations. Fig. 54 outlines the structures of some Pt-based photosensitizers that are highlighted here. Platinum complexes with conjugated acetylides are efficient phosphorescent chromophores. The photoredox and electrochemical properties of Pt acetylides can be readily modified by installing different substituents on the alkynyl ligands or by changing the auxiliary ligands. Eisenberg and coworkers reported a suite of terpyridyl Pt acetylide complexes (Pt1–Pt4) that exhibit photoluminesce with lmax between 595 and 625 nm from 3MLCT emissive states.259–261 These complexes were applied as photosensitizers for the reduction of water to generate H2, proceeding via the oxidative quenching of the photosensitizer’s excited state by methyl viologen, using triethylamine and colloidal Pt as the sacrificial reductant and catalyst, respectively. Bipyridine-supported diacetylide complex Pt5 and its detailed photophysical properties have been extensively explored.262,263 This complex is brightly emissive in fluid solution with d(Pt)!p (diimine) MLCT excited states. However, it displays weak absorbance of visible light and the excited-state lifetime is relatively low (t ¼ 1.3 ms). Castellano and coworkers incorporated pyrenyl substituents on the alkynyl ligands in place of phenyl groups to prepare Pt6, which luminesces from excited state centered on the pyrenylacetylides with triplet-state lifetime up to 48 ms.264

Fig. 54 Structures of some Pt-based organometallic photosensitizers.

Organometallic Photosensitizers

333

Che et al. reported a blue phosphor, Pt7, as a strong photoreductant indicated by excited-state oxidation potential [E(PtIII/ PtII)] of −2.63 V vs. Fc+/Fc.265 This complex features tetradentate bis(phenolate-NHC) ligands with N-carbazolyl motifs. The electron-donating N-carbazolyl groups and chelating effect improve the stability of the complex both in the ground and excited states, making it a robust photosensitizer in photocatalytic reactions. Choi and coworkers applied Pt(ppy)(acac) (Pt8, ppy ¼ 2-phenylpyridine, acac ¼ acetylacetonate) in photocatalytic trifluoromethylation of terminal alkenes. The excited-state redox properties of this complex can be controlled by changing the substituents on the ppy ligand.266 Complex Pt9, which contains a ppy ligand and a chelating diphosphine construct, shows intense absorption between 300 and 450 nm and luminesces at lmax ¼ 543 nm.267 Photocatalytic difluoroalkylation using complex Pt9 was carried out using cinnamic acid or alkyne substrates partnered with difluorohaloacetate substrates. These reactions form difluoroalkyl alkenes and difluoroalkyl alkenyl iodides, E-selective with cinnamic acid substrates and Z-selective with alkyne substrates. Even though group 10 complexes have been much less frequently used as photosensitizers compared to iridium(III), in many ways their synthetic chemistry and photochemistry are similar, particularly for Pt(II). Like Ir(III), Pt(II) and Pd(II) can be readily chelated with C^N and C^C cyclometalating ligands, with homoleptic and heteroleptic complexes both possible. As a result, the excited-state energy, lifetime, and redox potentials of these cyclometalated group 10 complexes can be tuned in much the same way as Ir(III) cyclometalates. Their lack of sustained development in photocatalysis is at least partly due to their inferior photostability, brought on by the four-coordinate planar structures of d8 complexes that open up many potential associative addition/substitution pathways during irradiation. That said, the lower coordination numbers and greater substitutional lability of Pt(II) and Pd(II) photosensitizers can be an advantage, in that it allows excited-state outer-sphere redox chemistry to be coupled with inner-sphere bond-activation chemistry. This feature was used in the example of Pd1 given above, which could mediate radical photoredox reactions and two-electron coupling reactions in the same reaction pot, and there are also recent examples where photochemical activation of cyclometalated platinum complexes can promote CdH activation, without accompanying outer-sphere redox chemistry.268,269 This unique ability of group 10 complexes to couple long-lived triplet excited states with inner-sphere activation of small-molecule and organic substrates may enable new transformations which are not possible with photoinduced electron-transfer or inner-sphere chemistry alone.

1.10.9

Coinage metal photosensitizers

There are many copper(I) complexes which have been developed as MLCT photosensitizers, normally with a homoleptic [Cu(diimine)2]+ or [Cu(diimine)(diphosphine)]+ structure.270–272 These copper photosensitizers are attractive alternatives to the precious-metal photosensitizers, and the d10 metal centers do not have the ligand-field states that are deleterious to other first-row transition metal photosensitizers. However, they are not described in detail here, as they are not organometallic. However, an organometallic heteroleptic copper photosensitizer has been developed (Fig. 55). Reiser et al. reported [Cu(dpp)(binc)]+ [dpp ¼ 2,9-diphenyl-1,10-phenanthroline, binc ¼ bis(2-isocyanophenyl)phenylphosphonate], Cu1, showing an excited state lifetime of 17 ms when immobilized in PMMA matrices [PMMA ¼ poly(methyl methacrylate)].273 The isocyanide ligands suppress structural reorganization upon excitation, which lengthens the lifetime and improves the photocatalytic performance for atom-transfer radical addition (ATRA) reactions and allylation of organohalides with allyltrimethylsilane. In recognition of the isolobal relationship between square-planar platinum(II) and gold(III) complexes, Che and coworkers synthesized a NHC-supported gold(III) complex (Au1) with a p-conjugated bis-cyclometalated ligand.274 This complex shows an exceptionally long excited state lifetime of 506 ms in dichloromethane solution and was demonstrated to be effective for both oxidative C(sp3)dH bond functionalization and for proton reduction to H2 upon white-light irradiation with [Co(dmgh)2(py)Cl] as the catalyst.

Fig. 55 Structures of coinage metal photosensitizers.

334

1.10.10

Organometallic Photosensitizers

Summary and conclusions

This article has introduced the physical principles of transition metal organometallic photosensitizers and highlighted the diverse structure types that have been developed over several decades of research. Recognizing that there is no “one size fits all” photosensitizer, it has been critical for the advancement of photocatalysis and photovoltaic research that a wide variety of photosensitizers have emerged, with a large range of photophysical and redox characteristics. The availability of these photosensitizers with diverse attributes means that for many applications it is possible to select a photosensitizer that is already ideally suited, and this is typically what is done in photoredox catalysis method development. For other emerging applications it is still necessary to design new photosensitizers and further tune the excited-state redox properties. For example, photoredox transformations on unactivated organohalide substrates remain a challenge, and visible-light activation of these molecules usually requires additional energy input (applied potential or two-photon) and/or wasteful secondary reagents. The continued design of potent visible-light photoreductants, akin to the group 6 isocyanide complexes and heteroleptic Ir(C^N)2(L^X) (L^X ¼ electron-rich chelate) complexes described in this article, will facilitate discovery of photoredox transformations on a much wider range of substrates. In addition, the lack of panchromatic absorbing photosensitizers outside of the well-studied ruthenium polypyridyl family represents a particularly significant fundamental challenge in photosensitizer design. One other future direction that will likely experience significant development in the near future is the coupling of excited-state electron or energy-transfer chemistry with inner-sphere bond activation. This is particularly feasible with photosensitizers from groups 10 and 11 that have coordination numbers 4, and it could make some existing transformations more efficient by reducing the need for secondary catalysts or sacrificial reagents, as well as open the door to some exciting new reactivity in the contexts of small-molecule activation for solar fuels chemistry and organic photoredox methodology. Given the direction of the organometallic field as a whole, we also envision the design of earth-abundant photosensitizers will be pursued with particular vigor, to be able to replace highly successful but highly scarce metals like iridium, platinum, and ruthenium, which have dominated the field so far. First-row metals pose their own set of unique challenges in photosensitizer design, in particular the low-lying ligand-field (d–d) excited states that can deactivate the desired charge-transfer state and lead to ligand-loss photochemical degradation once populated. For this reason copper(I) photosensitizers, which are d10 and thus do not have any ligand-field states, are especially promising as next-generation earth-abundant photosensitizers. Most existing copper(I) designs use some combination of chelating 1,10-phananthroline derivatives and diphosphine supporting ligands, and expansion to other ligand classes is certainly warranted. That said, given the relative scarcity of stable copper(I) complexes that are truly organometallic, i.e. having one or more metal-carbon bond, it seems unlikely that organocopper compounds will make a huge impact. Over the last several decades the depth of understanding of organometallic photosensitizers has greatly matured, but there are still opportunities for creative synthetic chemists to introduce new organometallic platforms that will teach us more insights into fundamental molecular photophysics and impact the many synthetic, catalytic, and device applications that rely on photosensitizer molecules.

Acknowledgments We acknowledge the National Science Foundation (grant number CHE-1846831) and the Welch Foundation (grant number E-1887) for supporting our research on organometallic photosensitizers.

References 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19.

Han, Z.; Eisenberg, R. Acc. Chem. Res. 2014, 47 (8), 2537–2544. Carella, A.; Borbone, F.; Centore, R. Front. Chem. 2018, 6. Prier, C. K.; Rankic, D. A.; MacMillan, D. W. C. Chem. Rev. 2013, 113 (7), 5322–5363. Tucker, J. W.; Stephenson, C. R. J. J. Org. Chem. 2012, 77 (4), 1617–1622. Choung, K. S.; Marroquin, K.; Teets, T. S. Chem. Sci. 2019, 10 (19), 5124–5132. Ma, D.-L.; Ma, V. P.-Y.; Chan, D. S.-H.; Leung, K.-H.; He, H.-Z.; Leung, C.-H. Coord. Chem. Rev. 2012, 256 (23), 3087–3113. Bonesi, S. M.; Mella, M.; d’Alessandro, N.; Aloisi, G. G.; Vanossi, M.; Albini, A. J. Org. Chem. 1998, 63 (26), 9946–9955. Wang, Z.-S.; Sayama, K.; Sugihara, H. J. Phys. Chem. B 2005, 109 (47), 22449–22455. Romero, N. A.; Nicewicz, D. A. Chem. Rev. 2016, 116 (17), 10075–10166. Sato, S.; Arai, T.; Morikawa, T.; Uemura, K.; Suzuki, T. M.; Tanaka, H.; Kajino, T. J. Am. Chem. Soc. 2011, 133 (39), 15240–15243. Hautala, R. R.; Little, J.; Sweet, E. Sol. Energy 1977, 19 (5), 503–508. Hockin, B. M.; Li, C.; Robertson, N.; Zysman-Colman, E. Cat. Sci. Technol. 2019, 9 (4), 889–915. Glaser, F.; Wenger, O. S. Coord. Chem. Rev. 2020, 405, 213129. Yin, J.-F.; Velayudham, M.; Bhattacharya, D.; Lin, H.-C.; Lu, K.-L. Coord. Chem. Rev. 2012, 256 (23–24), 3008–3035. Islam, A.; Sugihara, H.; Arakawa, H. J. Photochem. Photobiol. Chem. 2003, 158 (2–3), 131–138. Yoon, T. P.; Ischay, M. A.; Du, J. Nat. Chem. 2010, 2, 527–532. Gill, M. R.; Thomas, J. A. Chem. Soc. Rev. 2012, 41 (8), 3179. Vougioukalakis, G. C.; Philippopoulos, A. I.; Stergiopoulos, T.; Falaras, P. Coord. Chem. Rev. 2011, 255 (21–22), 2602–2621. Browne, W. R., Holder, A. A., Lawrence, M. A., Bullock, J. L., Lilge, L., Eds.; In Ruthenium Complexes: Photochemical and Biomedical Applications; Wiley-VCH Verlag GmbH & Co. KGaA: Weinheim, Germany, 2018. 20. Mishra, A.; Fischer, M. K. R.; Bäuerle, P. Angew. Chem. Int. Ed. 2009, 48 (14), 2474–2499.

Organometallic Photosensitizers 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. 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.

335

Wu, Y.; Zhu, W. Chem. Soc. Rev. 2013, 42 (5), 2039–2058. Yersin, H., Ed.; In Highly Efficient OLEDs With Phosphorescent Materials; Wiley-VCH: Weinheim, 2008. Yersin, H. Transition Metal and Rare Earth Compounds: Excited States, Transitions, Interactions III; Springer Berlin Heidelberg: Berlin, Heidelberg, 2004; pp 1–26. Zysman-Colman, E., Ed.; In Iridium(III) in Optoelectronic and Photonics Applications; John Wiley & Sons, Inc: Chichester, West Sussex, 2017. Wang, P.; Guo, S.; Wang, H.-J.; Chen, K.-K.; Zhang, N.; Zhang, Z.-M.; Lu, T.-B. Nat. Commun. 2019, 10 (1), 3155. Galletta, M.; Campagna, S.; Quesada, M.; Ulrich, G.; Ziessel, R. Chem. Commun. 2005, 33, 4222–4224. McCusker, C. E.; Castellano, F. N. Inorg. Chem. 2013, 52 (14), 8114–8120. Wächtler, M.; Kübel, J.; Barthelmes, K.; Winter, A.; Schmiedel, A.; Pascher, T.; Lambert, C.; Schubert, U. S.; Dietzek, B. Phys. Chem. Chem. Phys. 2016, 18 (4), 2350–2360. P. Wang, R. Dong, S. Guo, J. Zhao, Z.-M. Zhang and T.-B. Lu, Natl. Sci. Rev. 7 (9), 2020, 1459–1467. Teets, T. S.; Nocera, D. G. Chem. Commun. 2011, 47 (33), 9268–9274. Ishida, H. In Carbon Dioxide Chemistry, Capture and Oil Recovery; Karamé, I., Shaya, J., Srour, H., Eds.; InTech, 2018. Förster, T. Ann. Phys. 1948, 437 (1–2), 55–75. Jares-Erijman, E. A.; Jovin, T. M. Nat. Biotechnol. 2003, 21 (11), 1387–1395. Sapsford, K. E.; Berti, L.; Medintz, I. L. Angew. Chem. Int. Ed. 2006, 45 (28), 4562–4589. Strieth-Kalthoff, F.; James, M. J.; Teders, M.; Pitzer, L.; Glorius, F. Chem. Soc. Rev. 2018, 47 (19), 7190–7202. Bunt, G.; Wouters, F. S. Biophys. Rev. 2017, 9 (2), 119–129. Truong, K.; Ikura, M. Curr. Opin. Struct. Biol. 2001, 11 (5), 573–578. Kavarnos, G. J.; Turro, N. J. Chem. Rev. 1986, 86 (2), 401–449. Dexter, D. L. J. Chem. Phys. 1953, 21 (5), 836–850. Welin, E. R.; Le, C.; Arias-Rotondo, D. M.; McCusker, J. K.; MacMillan, D. W. C. Science 2017, 355 (6323), 380–385. Lu, Z.; Yoon, T. P. Angew. Chem. Int. Ed. 2012, 51 (41), 10329–10332. Zhang, Y.; Petersen, J. L.; Milsmann, C. Organometallics 2018, 37 (23), 4488–4499. Zhang, Y.; Lee, T. S.; Petersen, J. L.; Milsmann, C. J. Am. Chem. Soc. 2018, 140 (18), 5934–5947. Zhang, Y.; Petersen, J. L.; Milsmann, C. J. Am. Chem. Soc. 2016, 138 (40), 13115–13118. Mann, K. R.; Gray, H. B.; Hammond, G. S. J. Am. Chem. Soc. 1977, 99 (1), 306–307. Sattler, W.; Henling, L. M.; Winkler, J. R.; Gray, H. B. J. Am. Chem. Soc. 2015, 137 (3), 1198–1205. Sattler, W.; Ener, M. E.; Blakemore, J. D.; Rachford, A. A.; LaBeaume, P. J.; Thackeray, J. W.; Cameron, J. F.; Winkler, J. R.; Gray, H. B. J. Am. Chem. Soc. 2013, 135 (29), 10614–10617. Herr, P.; Glaser, F.; Büldt, L. A.; Larsen, C. B.; Wenger, O. S. J. Am. Chem. Soc. 2019, 141 (36), 14394–14402. Büldt, L. A.; Guo, X.; Prescimone, A.; Wenger, O. S. Angew. Chem. Int. Ed. 2016, 55 (37), 11247–11250. Büldt, L. A.; Guo, X.; Vogel, R.; Prescimone, A.; Wenger, O. S. J. Am. Chem. Soc. 2017, 139 (2), 985–992. Kjær, K. S.; Kaul, N.; Prakash, O.; Chábera, P.; Rosemann, N. W.; Honarfar, A.; Gordivska, O.; Fredin, L. A.; Bergquist, K.-E.; Häggström, L.; Ericsson, T.; Lindh, L.; Yartsev, A.; Styring, S.; Huang, P.; Uhlig, J.; Bendix, J.; Strand, D.; Sundström, V.; Persson, P.; Lomoth, R.; Wärnmark, K. Science 2019, 363 (6424), 249–253. Rosemann, N. W.; Chábera, P.; Prakash, O.; Kaufhold, S.; Wärnmark, K.; Yartsev, A.; Persson, P. J. Am. Chem. Soc. 2020, 142 (19), 8565–8569. Juris, A.; Balzani, V.; Barigelletti, F.; Campagna, S.; Belser, P.; von Zelewsky, A. Coord. Chem. Rev. 1988, 84, 85–277. Bessho, T.; Yoneda, E.; Yum, J.-H.; Guglielmi, M.; Tavernelli, I.; Imai, H.; Rothlisberger, U.; Nazeeruddin, M. K.; Grätzel, M. J. Am. Chem. Soc. 2009, 131 (16), 5930–5934. Bomben, P. G.; Robson, K. C. D.; Koivisto, B. D.; Berlinguette, C. P. Coord. Chem. Rev. 2012, 256 (15), 1438–1450. Bomben, P. G.; Koivisto, B. D.; Berlinguette, C. P. Inorg. Chem. 2010, 49 (11), 4960–4971. Wadman, S. H.; Kroon, J. M.; Bakker, K.; Havenith, R. W. A.; van Klink, G. P. M.; van Koten, G. Organometallics 2010, 29 (7), 1569–1579. Wadman, S. H.; Kroon, J. M.; Bakker, K.; Lutz, M.; Spek, A. L.; Klink, G. P. M.; Koten, G. Chem. Commun. 2007, (19), 1907–1909. Wadman, S. H.; Lutz, M.; Tooke, D. M.; Spek, A. L.; Hartl, F.; Havenith, R. W. A.; van Klink, G. P. M.; van Koten, G. Inorg. Chem. 2009, 48 (5), 1887–1900. Bomben, P. G.; Robson, K. C. D.; Sedach, P. A.; Berlinguette, C. P. Inorg. Chem. 2009, 48 (20), 9631–9643. Chang, W.-C.; Chen, H.-S.; Li, T.-Y.; Hsu, N.-M.; Tingare, Y. S.; Li, C.-Y.; Liu, Y.-C.; Su, C.; Li, W.-R. Angew. Chem. Int. Ed. 2010, 49 (44), 8161–8164. Son, S. U.; Park, K. H.; Lee, Y.-S.; Kim, B. Y.; Choi, C. H.; Lah, M. S.; Jang, Y. H.; Jang, D.-J.; Chung, Y. K. Inorg. Chem. 2004, 43 (22), 6896–6898. Torres, J.; Carrión, M. C.; Leal, J.; Castañeda, G.; Manzano, B. R.; Jalón, F. A. J. Organomet. Chem. 2019, 898, 120880. Nonoyama, M. Bull. Chem. Soc. Jpn. 1974, 47 (3), 767–768. Garces, F. O.; King, K. A.; Watts, R. J. Inorg. Chem. 1988, 27 (20), 3464–3471. Sprouse, S.; King, K. A.; Spellane, P. J.; Watts, R. J. J. Am. Chem. Soc. 1984, 106 (22), 6647–6653. King, K. A.; Spellane, P. J.; Watts, R. J. J. Am. Chem. Soc. 1985, 107 (5), 1431–1432. Dedeian, K.; Djurovich, P. I.; Garces, F. O.; Carlson, G.; Watts, R. J. Inorg. Chem. 1991, 30 (8), 1685–1687. Nguyen, J. D.; D’Amato, E. M.; Narayanam, J. M. R.; Stephenson, C. R. J. Nat. Chem. 2012, 4 (10), 854–859. McNally, A.; Prier, C. K.; MacMillan, D. W. C. Science 2011, 334 (6059), 1114–1117. Shih, H.-W.; Vander Wal, M. N.; Grange, R. L.; MacMillan, D. W. C. J. Am. Chem. Soc. 2010, 132 (39), 13600–13603. Xu, J.; Jung, K.; Atme, A.; Shanmugam, S.; Boyer, C. J. Am. Chem. Soc. 2014, 136 (14), 5508–5519. Yasu, Y.; Koike, T.; Akita, M. Angew. Chem. Int. Ed. 2012, 51 (38), 9567–9571. Jiang, H.; Cheng, Y.; Wang, R.; Zheng, M.; Zhang, Y.; Yu, S. Angew. Chem. Int. Ed. 2013, 52 (50), 13289–13292. Walker, M. M.; Koronkiewicz, B.; Chen, S.; Houk, K. N.; Mayer, J. M.; Ellman, J. A. J. Am. Chem. Soc. 2020, 142 (18), 8194–8202. Arora, A.; Weaver, J. D. Acc. Chem. Res. 2016, 49 (10), 2273–2283. Nacsa, E. D.; MacMillan, D. W. C. J. Am. Chem. Soc. 2018, 140 (9), 3322–3330. Singh, A.; Teegardin, K.; Kelly, M.; Prasad, K. S.; Krishnan, S.; Weaver, J. D. J. Organomet. Chem. 2015, 776, 51–59. Singh, A.; Fennell, C. J.; Weaver, J. D. Chem. Sci. 2016, 7 (11), 6796–6802. Kerzig, C.; Guo, X.; Wenger, O. S. J. Am. Chem. Soc. 2019, 141 (5), 2122–2127. Kerzig, C.; Wenger, O. S. Chem. Sci. 2019, 10 (48), 11023–11029. Pfund, B.; Steffen, D. M.; Schreier, M. R.; Bertrams, M.-S.; Ye, C.; Börjesson, K.; Wenger, O. S.; Kerzig, C. J. Am. Chem. Soc. 2020, 142 (23), 10468–10476. Tamayo, A. B.; Alleyne, B. D.; Djurovich, P. I.; Lamansky, S.; Tsyba, I.; Ho, N. N.; Bau, R.; Thompson, M. E. J. Am. Chem. Soc. 2003, 125 (24), 7377–7387. Maity, A.; Anderson, B. L.; Deligonul, N.; Gray, T. G. Chem. Sci. 2013, 4 (3), 1175–1181. Tinker, L. L.; McDaniel, N. D.; Cline, E. D.; Bernhard, S. Inorg. Synth. 2010, 35, 168–173. King, K. A.; Watts, R. J. J. Am. Chem. Soc. 1987, 109 (5), 1589–1590. Ohsawa, Y.; Sprouse, S.; King, K. A.; DeArmond, M. K.; Hanck, K. W.; Watts, R. J. J. Phys. Chem. 1987, 91, 1047–1054. Wilde, A. P.; King, K. A.; Watts, R. J. J. Phys. Chem. 1991, 95 (2), 629–634. Neve, F.; Crispini, A.; Campagna, S.; Serroni, S. Inorg. Chem. 1999, 38 (10), 2250–2258. Neve, F.; Crispini, A. Eur. J. Inorg. Chem. 2000, 2000 (5), 1039–1043. Neve, F.; Crispini, A.; Loiseau, F.; Campagna, S. J. Chem. Soc. Dalton Trans. 2000, (9), 1399–1401.

336 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. 131. 132. 133. 134. 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. 160. 161.

Organometallic Photosensitizers Glusac, K. D.; Jiang, S.; Schanze, K. S. Chem. Commun. 2002, (21), 2504–2505. Cunningham, G. B.; Li, Y.; Liu, S.; Schanze, K. S. J. Phys. Chem. B 2003, 107 (46), 12569–12572. Griffiths, P. M.; Loiseau, F.; Puntoriero, F.; Serroni, S.; Campagna, S. Chem. Commun. 2000, (23), 2297–2298. Neve, F.; Crispini, A.; Serroni, S.; Loiseau, F.; Campagna, S. Inorg. Chem. 2001, 40 (6), 1093–1101. Lo, K. K.-W.; Ng, D. C.-M.; Chung, C.-K. Organometallics 2001, 20 (24), 4999–5001. Lo, K. K.-W.; Chung, C.-K.; Zhu, N. Chem. A Eur. J. 2003, 9 (2), 475–483. Lo, K. K.-W.; Chung, C.-K.; Lee, T. K.-M.; Lui, L.-H.; Tsang, K. H.-K.; Zhu, N. Inorg. Chem. 2003, 42 (21), 6886–6897. Lo, K. K.-W.; Chan, J. S.-W.; Chung, C.-K.; Tsang, V. W.-H.; Zhu, N. Inorg. Chim. Acta 2004, 357 (10), 3109–3118. Lo, K. K.-W.; Chan, J. S.-W.; Lui, L.-H.; Chung, C.-K. Organometallics 2004, 23 (13), 3108–3116. Slinker, J. D.; Gorodetsky, A. A.; Lowry, M. S.; Wang, J.; Parker, S.; Rohl, R.; Bernhard, S.; Malliaras, G. G. J. Am. Chem. Soc. 2004, 126 (9), 2763–2767. Lowry, M. S.; Hudson, W. R.; Pascal, R. A.; Bernhard, S. J. Am. Chem. Soc. 2004, 126 (43), 14129–14135. Goldsmith, J. I.; Hudson, W. R.; Lowry, M. S.; Anderson, T. H.; Bernhard, S. J. Am. Chem. Soc. 2005, 127 (20), 7502–7510. Cline, E. D.; Adamson, S. E.; Bernhard, S. Inorg. Chem. 2008, 47 (22), 10378–10388. Lowry, M. S.; Goldsmith, J. I.; Slinker, J. D.; Rohl, R.; Pascal, R. A.; Malliaras, G. G.; Bernhard, S. Chem. Mater. 2005, 17 (23), 5712–5719. Nagib, D. A.; Scott, M. E.; MacMillan, D. W. C. J. Am. Chem. Soc. 2009, 131 (31), 10875–10877. Wallentin, C.-J.; Nguyen, J. D.; Finkbeiner, P.; Stephenson, C. R. J. J. Am. Chem. Soc. 2012, 134 (21), 8875–8884. Jin, J.; MacMillan, D. W. C. Nature 2015, 525 (7567), 87–90. Zhu, S.; Das, A.; Bui, L.; Zhou, H.; Curran, D. P.; Rueping, M. J. Am. Chem. Soc. 2013, 135 (5), 1823–1829. Noble, A.; McCarver, S. J.; MacMillan, D. W. C. J. Am. Chem. Soc. 2015, 137 (2), 624–627. Koike, T.; Akita, M. Acc. Chem. Res. 2016, 49 (9), 1937–1945. Nguyen, J. D.; Matsuura, B. S.; Stephenson, C. R. J. J. Am. Chem. Soc. 2014, 136 (4), 1218–1221. Rono, L. J.; Yayla, H. G.; Wang, D. Y.; Armstrong, M. F.; Knowles, R. R. J. Am. Chem. Soc. 2013, 135 (47), 17735–17738. Jia, J.; Kancherla, R.; Rueping, M.; Huang, L. Chem. Sci. 2020, 11 (19), 4954–4959. Simons, R. T.; Scott, G. E.; Kanegusuku, A. G.; Roizen, J. L. J. Org. Chem. 2020, 85 (10), 6380–6391. Schwarz, J. L.; Kleinmans, R.; Paulisch, T. O.; Glorius, F. J. Am. Chem. Soc. 2020, 142 (5), 2168–2174. Li, J.; Chen, J.; Sang, R.; Ham, W.-S.; Plutschack, M. B.; Berger, F.; Chabbra, S.; Schnegg, A.; Genicot, C.; Ritter, T. Nat. Chem. 2020, 12 (1), 56–62. Thullen, S. M.; Treacy, S. M.; Rovis, T. J. Am. Chem. Soc. 2019, 141 (36), 14062–14067. Ye, J.; Kalvet, I.; Schoenebeck, F.; Rovis, T. Nat. Chem. 2018, 10 (10), 1037–1041. Wang, R.; Ma, M.; Gong, X.; Panetti, G. B.; Fan, X.; Walsh, P. J. Org. Lett. 2018, 20 (8), 2433–2436. Chen, D.-F.; Chu, J. C. K.; Rovis, T. J. Am. Chem. Soc. 2017, 139 (42), 14897–14900. Deng, Y.; Liu, Q.; Smith, A. B. J. Am. Chem. Soc. 2017, 139 (28), 9487–9490. Nielsen, M. K.; Shields, B. J.; Liu, J.; Williams, M. J.; Zacuto, M. J.; Doyle, A. G. Angew. Chem. Int. Ed. 2017, 56 (25), 7191–7194. Paul, A.; Smith, M. D.; Vannucci, A. K. J. Org. Chem. 2017, 82 (4), 1996–2003. Yayla, H. G.; Wang, H.; Tarantino, K. T.; Orbe, H. S.; Knowles, R. R. J. Am. Chem. Soc. 2016, 138 (34), 10794–10797. Scholz, S. O.; Farney, E. P.; Kim, S.; Bates, D. M.; Yoon, T. P. Angew. Chem. Int. Ed. 2016, 55 (6), 2239–2242. Musacchio, A. J.; Nguyen, L. Q.; Beard, G. H.; Knowles, R. R. J. Am. Chem. Soc. 2014, 136 (35), 12217–12220. Tellis, J. C.; Primer, D. N.; Molander, G. A. Science 2014, 345 (6195), 433–436. Primer, D. N.; Karakaya, I.; Tellis, J. C.; Molander, G. A. J. Am. Chem. Soc. 2015, 137 (6), 2195–2198. Heitz, D. R.; Tellis, J. C.; Molander, G. A. J. Am. Chem. Soc. 2016, 138 (39), 12715–12718. Choi, G. J.; Zhu, Q.; Miller, D. C.; Gu, C. J.; Knowles, R. R. Nature 2016, 539 (7628), 268–271. Mayo, E. I.; Kils, K.; Tirrell, T.; Djurovich, P. I.; Tamayo, A.; Thompson, M. E.; Lewis, N. S.; Gray, H. B. Photochem. Photobiol. Sci. 2006, 5 (10), 871. Waern, J. B.; Desmarets, C.; Chamoreau, L.-M.; Amouri, H.; Barbieri, A.; Sabatini, C.; Ventura, B.; Barigelletti, F. Inorg. Chem. 2008, 47 (8), 3340–3348. Dragonetti, C.; Valore, A.; Colombo, A.; Righetto, S.; Trifiletti, V. Inorg. Chim. Acta 2012, 388, 163–167. Ning, Z.; Zhang, Q.; Wu, W.; Tian, H. J. Organomet. Chem. 2009, 694 (17), 2705–2711. Hagfeldt, A.; Boschloo, G.; Sun, L.; Kloo, L.; Pettersson, H. Chem. Rev. 2010, 110 (11), 6595–6663. Sinopoli, A.; Black, F. A.; Wood, C. J.; Gibson, E. A.; Elliott, P. I. P. Dalton Trans. 2017, 46 (5), 1520–1530. Li, L.; Zhang, S.; Xu, L.; Han, L.; Chen, Z.-N.; Luo, J. Inorg. Chem. 2013, 52 (21), 12323–12325. Li, L.; Zhang, S.; Xu, L.; Wang, J.; Shi, L.-X.; Chen, Z.-N.; Hong, M.; Luo, J. Chem. Sci. 2014, 5 (10), 3808. Li, L.-P.; Ye, B.-H. Inorg. Chem. 2019, 58 (12), 7775–7784. Fihri, A.; Artero, V.; Pereira, A.; Fontecave, M. Dalton Trans. 2008, 41, 5567. Lentz, C.; Schott, O.; Auvray, T.; Hanan, G.; Elias, B. Inorg. Chem. 2017, 56 (18), 10875–10881. Lentz, C.; Schott, O.; Auvray, T.; Hanan, G. S.; Elias, B. Dalton Trans. 2019, 48 (41), 15567–15576. Gärtner, F.; Cozzula, D.; Losse, S.; Boddien, A.; Anilkumar, G.; Junge, H.; Schulz, T.; Marquet, N.; Spannenberg, A.; Gladiali, S.; Beller, M. Chem. A Eur. J. 2011, 17 (25), 6998–7006. Polo, A. S.; Itokazu, M. K.; Murakami Iha, N. Y. Coord. Chem. Rev. 2004, 248, 1343–1361. Bevernaegie, R.; Marcélis, L.; Laramée-Milette, B.; De Winter, J.; Robeyns, K.; Gerbaux, P.; Hanan, G. S.; Elias, B. Inorg. Chem. 2018, 57 (3), 1356–1367. Bevernaegie, R.; Wehlin, S. A. M.; Piechota, E. J.; Abraham, M.; Philouze, C.; Meyer, G. J.; Elias, B.; Troian-Gautier, L. J. Am. Chem. Soc. 2020, 142 (6), 2732–2737. Hasan, K.; Zysman-Colman, E. Inorg. Chem. 2012, 51 (22), 12560–12564. Nazeeruddin, M. K.; Kay, A.; Rodicio, I.; Humphry-Baker, R.; Mueller, E.; Liska, P.; Vlachopoulos, N.; Graetzel, M. J. Am. Chem. Soc. 1993, 115 (14), 6382–6390. Hasan, K.; Zysman-Colman, E. Eur. J. Inorg. Chem. 2013, 2013 (25), 4421–4429. Hasan, K.; Wang, J.; Pal, A. K.; Hierlinger, C.; Guerchais, V.; Sen Soo, H.; García, F.; Zysman-Colman, E. Sci. Rep. 2017, 7 (1), 15520. Sinopoli, A.; Wood, C. J.; Gibson, E. A.; Elliott, P. I. P. Dyes Pigments 2017, 140, 269–277. Gennari, M.; Légalité, F.; Zhang, L.; Pellegrin, Y.; Blart, E.; Fortage, J.; Brown, A. M.; Deronzier, A.; Collomb, M.-N.; Boujtita, M.; Jacquemin, D.; Hammarström, L.; Odobel, F. J. Phys. Chem. Lett. 2014, 5 (13), 2254–2258. Henwood, A. F.; Hu, Y.; Sajjad, M. T.; Thalluri, G.; Ghosh, S. S.; Cordes, D. B.; Slawin, A. M. Z.; Samuel, I. D. W.; Robertson, N.; Zysman-Colman, E. Chem. A Eur. J. 2015, 21 (52), 19128–19135. Yuan, Y.-J.; Zhang, J.-Y.; Yu, Z.-T.; Feng, J.-Y.; Luo, W.-J.; Ye, J.-H.; Zou, Z.-G. Inorg. Chem. 2012, 51 (7), 4123–4133. Takizawa, S.; Shimada, K.; Sato, Y.; Murata, S. Inorg. Chem. 2014, 53 (6), 2983–2995. Takizawa, S.; Pérez-Bolívar, C.; Anzenbacher, P.; Murata, S. Eur. J. Inorg. Chem. 2012, 2012 (25), 3975–3979. Takizawa, S.; Ikuta, N.; Zeng, F.; Komaru, S.; Sebata, S.; Murata, S. Inorg. Chem. 2016, 55 (17), 8723–8735. Yarnell, J. E.; De La Torre, P.; Castellano, F. N. Eur. J. Inorg. Chem. 2017, 2017 (44), 5238–5245. Yang, M.; Yarnell, J. E.; El Roz, K.; Castellano, F. N. ACS Appl. Energy Mater. 2020, 3 (2), 1842–1853. Hallett, A. J.; White, N.; Wu, W.; Cui, X.; Horton, P. N.; Coles, S. J.; Zhao, J.; Pope, S. J. A. Chem. Commun. 2012, 48 (88), 10838.

Organometallic Photosensitizers

337

162. Lamansky, S.; Djurovich, P.; Murphy, D.; Abdel-Razzaq, F.; Lee, H.-E.; Adachi, C.; Burrows, P. E.; Forrest, S. R.; Thompson, M. E. J. Am. Chem. Soc. 2001, 123 (18), 4304–4312. 163. Lamansky, S.; Djurovich, P.; Murphy, D.; Abdel-Razzaq, F.; Kwong, R.; Tsyba, I.; Bortz, M.; Mui, B.; Bau, R.; Thompson, M. E. Inorg. Chem. 2001, 40 (7), 1704–1711. 164. Gao, R.; Ho, D. G.; Hernandez, B.; Selke, M.; Murphy, D.; Djurovich, P. I.; Thompson, M. E. J. Am. Chem. Soc. 2002, 124 (50), 14828–14829. 165. Djurovich, P. I.; Murphy, D.; Thompson, M. E.; Hernandez, B.; Gao, R.; Hunt, P. L.; Selke, M. Dalton Trans. 2007, 3763–3770. 166. Telitel, S.; Dumur, F.; Lepeltier, M.; Gigmes, D.; Fouassier, J.-P.; Lalevée, J. C. R. Chim. 2016, 19 (1–2), 71–78. 167. Tehfe, M.-A.; Lepeltier, M.; Dumur, F.; Gigmes, D.; Fouassier, J.-P.; Lalevée, J. Macromol. Chem. Phys. 2017, 218 (18), 1700192. 168. Wu, Z.; Liao, X.; Yuan, L.; Wang, Y.; Zheng, Y.; Zuo, J.; Pan, Y. Chem. A Eur. J. 2020, 26 (25), 5694–5700. 169. Baranoff, E.; Yum, J.-H.; Jung, I.; Vulcano, R.; Grätzel, M.; Nazeeruddin, M. K. Chem. Asian J. 2010, 5 (3), 496–499. 170. Huang, J.; Yu, J.; Guan, Z.; Jiang, Y. Appl. Phys. Lett. 2010, 97 (14), 143301. 171. Yu, J.; Zang, Y.; Li, H.; Huang, J. Thin Solid Films 2012, 520 (21), 6653–6657. 172. Yang, D.; Li, W.; Chu, B.; Su, Z.; Wang, J.; Zhang, G.; Zhang, F. Appl. Phys. Lett. 2011, 99 (19), 193301. 173. Zhen, H.; Hou, Q.; Li, K.; Ma, Z.; Fabiano, S.; Gao, F.; Zhang, F. J. Mater. Chem. A 2014, 2 (31), 12390. 174. Tan, G.; Liu, P.; Wu, H.; Yiu, S.-C.; Dai, F.; Feng, Y.-H.; Liu, X.; Qiu, Y.; Lo, Y. H.; Ho, C.-L.; Wong, W.-Y. J. Organomet. Chem. 2016, 812, 280–286. 175. Fan, P.; Zheng, Y.; Zheng, D.; Yu, J. Mater. Lett. 2017, 186, 161–164. 176. Qian, M.; Zhang, R.; Hao, J.; Zhang, W.; Zhang, Q.; Wang, J.; Tao, Y.; Chen, S.; Fang, J.; Huang, W. Adv. Mater. 2015, 27 (23), 3546–3552. 177. You, Y.; Park, S. Y. J. Am. Chem. Soc. 2005, 127 (36), 12438–12439. 178. Li, C.-X.; Tu, D.-S.; Yao, R.; Yan, H.; Lu, C.-S. Org. Lett. 2016, 18 (19), 4928–4931. 179. Zhou, H.; Lu, P.; Gu, X.; Li, P. Org. Lett. 2013, 15 (22), 5646–5649. 180. Wang, D.; Wu, Y.; Dong, H.; Qin, Z.; Zhao, D.; Yu, Y.; Zhou, G.; Jiao, B.; Wu, Z.; Gao, M.; Wang, G. Org. Electron. 2013, 14 (12), 3297–3305. 181. Lee, W.; Kwon, T.-H.; Kwon, J.; Kim, J.; Lee, C.; Hong, J.-I. New J. Chem. 2011, 35 (11), 2557. 182. Yun, M. H.; Lee, E.; Lee, W.; Choi, H.; Lee, B. R.; Song, M. H.; Hong, J.-I.; Kwon, T.-H.; Kim, J. Y. J. Mater. Chem. C 2014, 2 (47), 10195–10200. 183. Matt, B.; Moussa, J.; Chamoreau, L.-M.; Afonso, C.; Proust, A.; Amouri, H.; Izzet, G. Organometallics 2012, 31 (1), 35–38. 184. Matt, B.; Xiang, X.; Kaledin, A. L.; Han, N.; Moussa, J.; Amouri, H.; Alves, S.; Hill, C. L.; Lian, T.; Musaev, D. G.; Izzet, G.; Proust, A. Chem. Sci. 2013, 4 (4), 1737. 185. Matt, B.; Fize, J.; Moussa, J.; Amouri, H.; Pereira, A.; Artero, V.; Izzet, G.; Proust, A. Energ. Environ. Sci. 2013, 6 (5), 1504. 186. Liu, Y.; Ye, K.; Fan, Y.; Song, W.; Wang, Y.; Hou, Z. Chem. Commun. 2009, (25), 3699–3701. 187. Peng, T.; Bi, H.; Liu, Y.; Fan, Y.; Gao, H.; Wang, Y.; Hou, Z. J. Mater. Chem. 2009, 19 (43), 8072. 188. Rai, V. K.; Nishiura, M.; Takimoto, M.; Hou, Z. J. Mater. Chem. C 2014, 2 (27), 5317–5326. 189. Sahin, C.; Goren, A.; Varlikli, C. J. Organomet. Chem. 2014, 772–773, 68–78. 190. Rai, V. K.; Nishiura, M.; Takimoto, M.; Zhao, S.; Liu, Y.; Hou, Z. Inorg. Chem. 2012, 51 (2), 822–835. 191. Rai, V. K.; Nishiura, M.; Takimoto, M.; Hou, Z. J. Mater. Chem. C 2013, 1 (4), 677–689. 192. Yang, W.; Fu, H.; Song, Q.; Zhang, M.; Ding, Y. Organometallics 2011, 30 (1), 77–83. 193. Lai, P.-N.; Brysacz, C. H.; Alam, M. K.; Ayoub, N. A.; Gray, T. G.; Bao, J.; Teets, T. S. J. Am. Chem. Soc. 2018, 140 (32), 10198–10207. 194. Lai, P.-N.; Teets, T. S. J. Coord. Chem. 2019, 72 (8), 1238–1252. 195. Lai, P.; Teets, T. S. Chem. A Eur. J. 2019, 25 (23), 6026–6037. 196. Kabir, E.; Wu, Y.; Sittel, S.; Nguyen, B.-L.; Teets, T. S. Inorg. Chem. Front. 2020, 7 (6), 1362–1373. 197. Yu, Z.-T.; Yuan, Y.-J.; Cai, J.-G.; Zou, Z.-G. Chem. A Eur. J. 2013, 19 (4), 1303–1310. 198. Radwan, Y. K.; Maity, A.; Teets, T. S. Inorg. Chem. 2015, 54 (14), 7122–7131. 199. Maya, R. M.; Maity, A.; Teets, T. S. Organometallics 2016, 35 (17), 2890–2899. 200. Shon, J.-H.; Teets, T. S. Inorg. Chem. 2017, 56 (24), 15295–15303. 201. Shon, J.-H.; Sittel, S.; Teets, T. S. ACS Catal. 2019, 9 (9), 8646–8658. 202. Holmes, R. J.; Forrest, S. R.; Sajoto, T.; Tamayo, A.; Djurovich, P. I.; Thompson, M. E.; Brooks, J.; Tung, Y.-J.; D’Andrade, B. W.; Weaver, M. S.; Kwong, R. C.; Brown, J. J. Appl. Phys. Lett. 2005, 87, 243507. 203. Sajoto, T.; Djurovich, P. I.; Tamayo, A.; Yousufuddin, M.; Bau, R.; Thompson, M. E.; Holmes, R. J.; Forrest, S. R. Inorg. Chem. 2005, 44 (22), 7992–8003. 204. Lee, J.; Chen, H.-F.; Batagoda, T.; Coburn, C.; Djurovich, P. I.; Thompson, M. E.; Forrest, S. R. Nat. Mater. 2016, 15 (1), 92–98. 205. Pal, A. K.; Krotkus, S.; Fontani, M.; Mackenzie, C. F. R.; Cordes, D. B.; Slawin, A. M. Z.; Samuel, I. D. W.; Zysman-Colman, E. Adv. Mater. 2018, 30 (50), 1804231. 206. Na, H.; Cañada, L. M.; Wen, Z.; I-Chia Wu, J.; Teets, T. S. Chem. Sci. 2019, 10 (25), 6254–6260. 207. Chan, S. L.-F.; Lam, T. L.; Yang, C.; Yan, S.-C.; Cheng, N. M. Chem. Commun. 2015, 51 (37), 7799–7801. 208. Jin, J.; Shin, H.-W.; Park, J. H.; Park, J. H.; Kim, E.; Ahn, T. K.; Ryu, D. H.; Son, S. U. Organometallics 2013, 32 (14), 3954–3959. 209. Torres, J.; Carrión, M. C.; Leal, J.; Jalón, F. A.; Cuevas, J. V.; Rodríguez, A. M.; Castañeda, G.; Manzano, B. R. Inorg. Chem. 2018, 57 (3), 970–984. 210. Kessler, F.; Costa, R. D.; Di Censo, D.; Scopelliti, R.; Ortí, E.; Bolink, H. J.; Meier, S.; Sarfert, W.; Grätzel, M.; Nazeeruddin, M. K.; Baranoff, E. Dalton Trans. 2012, 41 (1), 180–191. 211. Monti, F.; Kessler, F.; Delgado, M.; Frey, J.; Bazzanini, F.; Accorsi, G.; Armaroli, N.; Bolink, H. J.; Ortí, E.; Scopelliti, R.; Nazeeruddin, M. K.; Baranoff, E. Inorg. Chem. 2013, 52 (18), 10292–10305. 212. Lam, T. L.; Lai, J.; Annapureddy, R. R.; Xue, M.; Yang, C.; Guan, Y.; Zhou, P.; Chan, S. L.-F. Inorg. Chem. 2017, 56 (18), 10835–10839. 213. Liu, B.; Monro, S.; Jabed, M. A.; Cameron, C. G.; Colón, K. L.; Xu, W.; Kilina, S.; McFarland, S. A.; Sun, W. Photochem. Photobiol. Sci. 2019, 18 (10), 2381–2396. 214. Amador, A. G.; Yoon, T. P. Angew. Chem. Int. Ed. 2016, 55 (7), 2304–2306. 215. Huo, H.; Shen, X.; Wang, C.; Zhang, L.; Röse, P.; Chen, L.-A.; Harms, K.; Marsch, M.; Hilt, G.; Meggers, E. Nature 2014, 515 (7525), 100–103. 216. Huo, H.; Fu, C.; Harms, K.; Meggers, E. J. Am. Chem. Soc. 2014, 136 (8), 2990–2993. 217. Huo, H.; Wang, C.; Harms, K.; Meggers, E. J. Am. Chem. Soc. 2015, 137 (30), 9551–9554. 218. Huo, H.; Huang, X.; Shen, X.; Harms, K.; Meggers, E. Synlett 2015, 27 (05), 749–753. 219. Wang, C.; Qin, J.; Shen, X.; Riedel, R.; Harms, K.; Meggers, E. Angew. Chem. Int. Ed. 2016, 55 (2), 685–688. 220. Zhang, X.; Qin, J.; Huang, X.; Meggers, E. Org. Chem. Front. 2018, 5 (2), 166–170. 221. Skubi, K. L.; Kidd, J. B.; Jung, H.; Guzei, I. A.; Baik, M.-H.; Yoon, T. P. J. Am. Chem. Soc. 2017, 139 (47), 17186–17192. 222. Zheng, J.; Swords, W. B.; Jung, H.; Skubi, K. L.; Kidd, J. B.; Meyer, G. J.; Baik, M.-H.; Yoon, T. P. J. Am. Chem. Soc. 2019, 141 (34), 13625–13634. 223. Manguin, R.; Pichon, D.; Tarrieu, R.; Vives, T.; Roisnel, T.; Dorcet, V.; Crévisy, C.; Miqueu, K.; Favereau, L.; Crassous, J.; Mauduit, M.; Baslé, O. Chem. Commun. 2019, 55 (43), 6058–6061. 224. Li, L.-P.; Peng, H.-L.; Wei, L.-Q.; Ye, B.-H. Inorg. Chem. 2019, 58 (1), 785–793. 225. Li, L.-P.; Yao, S.-Y.; Ou, Y.-L.; Wei, L.-Q.; Ye, B.-H. Organometallics 2017, 36 (17), 3257–3265. 226. Nazeeruddin, M. K.; Humphry-Baker, R.; Berner, D.; Rivier, S.; Zuppiroli, L.; Graetzel, M. J. Am. Chem. Soc. 2003, 125 (29), 8790–8797. 227. Shavaleev, N. M.; Monti, F.; Costa, R. D.; Scopelliti, R.; Bolink, H. J.; Ortí, E.; Accorsi, G.; Armaroli, N.; Baranoff, E.; Grätzel, M.; Nazeeruddin, M. K. Inorg. Chem. 2012, 51 (4), 2263–2271. 228. Shavaleev, N. M.; Monti, F.; Scopelliti, R.; Baschieri, A.; Sambri, L.; Armaroli, N.; Grätzel, M.; Nazeeruddin, M. K. Organometallics 2013, 32 (2), 460–467.

338 229. 230. 231. 232. 233. 234. 235. 236. 237. 238. 239. 240. 241. 242. 243. 244. 245. 246. 247. 248. 249. 250. 251. 252. 253. 254. 255. 256. 257. 258. 259. 260. 261. 262. 263. 264. 265. 266. 267. 268. 269. 270. 271. 272. 273. 274.

Organometallic Photosensitizers Shavaleev, N. M.; Monti, F.; Scopelliti, R.; Armaroli, N.; Grätzel, M.; Nazeeruddin, M. K. Organometallics 2012, 31 (17), 6288–6296. Maity, A.; Le, L. Q.; Zhu, Z.; Bao, J.; Teets, T. S. Inorg. Chem. 2016, 55, 2299–2308. Na, H.; Maity, A.; Teets, T. S. Dalton Trans. 2017, 46 (15), 5008–5016. Cañada, L. M.; Kölling, J.; Teets, T. S. Polyhedron 2020, 178, 114332. Favale, J. M.; Hauke, C. E.; Danilov, E. O.; Yarnell, J. E.; Castellano, F. N. Dalton Trans. 2020, 49 (29), 9995–10002. Abdel-Shafi, A. A.; Bourdelande, J. L.; Ali, S. S. Dalton Trans. 2007, 2510. Mamo, A.; Stefio, I.; Parisi, M. F.; Credi, A.; Venturi, M.; Di Pietro, C.; Campagna, S. Inorg. Chem. 1997, 36 (25), 5947–5950. Polson, M.; Fracasso, S.; Bertolasi, V.; Ravaglia, M.; Scandola, F. Inorg. Chem. 2004, 43 (6), 1950–1956. Polson, M.; Ravaglia, M.; Fracasso, S.; Garavelli, M.; Scandola, F. Inorg. Chem. 2005, 44 (5), 1282–1289. Whittle, V. L.; Williams, J. A. G. Inorg. Chem. 2008, 47 (15), 6596–6607. Wilkinson, A. J.; Puschmann, H.; Howard, J. A. K.; Foster, C. E.; Williams, J. A. G. Inorg. Chem. 2006, 45 (21), 8685–8699. Bexon, A. J. S.; Williams, J. A. G. C. R. Chim. 2005, 8 (8), 1326–1335. Williams, J. A. G.; Wilkinson, A. J.; Whittle, V. L. Dalton Trans. 2008, 16, 2081. Chirdon, D. N.; Transue, W. J.; Kagalwala, H. N.; Kaur, A.; Maurer, A. B.; Pintauer, T.; Bernhard, S. Inorg. Chem. 2014, 53 (3), 1487–1499. Porras, J. A.; Mills, I. N.; Transue, W. J.; Bernhard, S. J. Am. Chem. Soc. 2016, 138 (30), 9460–9472. Sato, S.; Morikawa, T.; Kajino, T.; Ishitani, O. Angew. Chem. Int. Ed. 2013, 52 (3), 988–992. Genoni, A.; Chirdon, D. N.; Boniolo, M.; Sartorel, A.; Bernhard, S.; Bonchio, M. ACS Catal. 2017, 7 (1), 154–160. Huang, H.; Banerjee, S.; Qiu, K.; Zhang, P.; Blacque, O.; Malcomson, T.; Paterson, M. J.; Clarkson, G. J.; Staniforth, M.; Stavros, V. G.; Gasser, G.; Chao, H.; Sadler, P. J. Nat. Chem. 2019, 11 (11), 1041–1048. Shinpuku, Y.; Inui, F.; Nakai, M.; Nakabayashi, Y. J. Photochem. Photobiol. Chem. 2011, 222 (1), 203–209. Wang, C.; Chen, L.-A.; Huo, H.; Shen, X.; Harms, K.; Gong, L.; Meggers, E. Chem. Sci. 2015, 6 (2), 1094–1100. Huang, X.; Meggers, E. Acc. Chem. Res. 2019, 52 (3), 833–847. Huang, X.; Li, X.; Xie, X.; Harms, K.; Riedel, R.; Meggers, E. Nat. Commun. 2017, 8 (1), 2245. Ma, J.; Harms, K.; Meggers, E. Chem. Commun. 2016, 52 (66), 10183–10186. Shen, X.; Harms, K.; Marsch, M.; Meggers, E. Chem. A Eur. J. 2016, 22 (27), 9102–9105. Steinlandt, P. S.; Zuo, W.; Harms, K.; Meggers, E. Chem. A Eur. J. 2019, 25 (67), 15333–15340. Grell, Y.; Hong, Y.; Huang, X.; Mochizuki, T.; Xie, X.; Harms, K.; Meggers, E. Organometallics 2019, 38 (20), 3948–3954. Tian, Y.-M.; Guo, X.-N.; Kuntze-Fechner, M. W.; Krummenacher, I.; Braunschweig, H.; Radius, U.; Steffen, A.; Marder, T. B. J. Am. Chem. Soc. 2018. Sieck, C.; Tay, M. G.; Thibault, M.-H.; Edkins, R. M.; Costuas, K.; Halet, J.-F.; Batsanov, A. S.; Haehnel, M.; Edkins, K.; Lorbach, A.; Steffen, A.; Marder, T. B. Chem. A Eur. J. 2016, 22 (30), 10523–10532. Hsu, Y.-C.; Wang, V. C.-C.; Au-Yeung, K.-C.; Tsai, C.-Y.; Chang, C.-C.; Lin, B.-C.; Chan, Y.-T.; Hsu, C.-P.; Yap, G. P. A.; Jurca, T.; Ong, T.-G. Angew. Chem. Int. Ed. 2018, 57 (17), 4622–4626. Chow, P. K.; Ma, C.; To, W.-P.; Tong, G. S. M.; Lai, S.-L.; Kui, S. C. F.; Kwok, W.-M.; Che, C.-M. Angew. Chem. Int. Ed. 2013, 52 (45), 11775–11779. Du, P.; Schneider, J.; Jarosz, P.; Eisenberg, R. J. Am. Chem. Soc. 2006, 128 (24), 7726–7727. Du, P.; Knowles, K.; Eisenberg, R. J. Am. Chem. Soc. 2008, 130 (38), 12576–12577. Du, P.; Schneider, J.; Jarosz, P.; Zhang, J.; Brennessel, W. W.; Eisenberg, R. J. Phys. Chem. B 2007, 111 (24), 6887–6894. Hissler, M.; Connick, W. B.; Geiger, D. K.; McGarrah, J. E.; Lipa, D.; Lachicotte, R. J.; Eisenberg, R. Inorg. Chem. 2000, 39 (3), 447–457. Hissler, M.; McGarrah, J. E.; Connick, W. B.; Geiger, D. K.; Cummings, S. D.; Eisenberg, R. Coord. Chem. Rev. 2000, 208 (1), 115–137. Pomestchenko, I. E.; Luman, C. R.; Hissler, M.; Ziessel, R.; Castellano, F. N. Inorg. Chem. 2003, 42 (5), 1394–1396. Li, K.; Wan, Q.; Yang, C.; Chang, X.-Y.; Low, K.-H.; Che, C.-M. Angew. Chem. Int. Ed. 2018, 57 (43), 14129–14133. Choi, W. J.; Choi, S.; Ohkubo, K.; Fukuzumi, S.; Cho, E. J.; You, Y. Chem. Sci. 2015, 6 (2), 1454–1464. Zhong, J.-J.; Yang, C.; Chang, X.-Y.; Zou, C.; Lu, W.; Che, C.-M. Chem. Commun. 2017, 53 (64), 8948–8951. Juliá, F.; González-Herrero, P. J. Am. Chem. Soc. 2016, 138 (16), 5276–5282. Poveda, D.; Vivancos, Á.; Bautista, D.; González-Herrero, P. Chem. Sci. 2020, 11 (44), 12095–12102. Heberle, M.; Tschierlei, S.; Rockstroh, N.; Ringenberg, M.; Frey, W.; Junge, H.; Beller, M.; Lochbrunner, S.; Karnahl, M. Chem. A Eur. J. 2017, 23 (2), 312–319. Housecroft, C. E.; Constable, E. C. Chem. Soc. Rev. 2015, 44 (23), 8386–8398. Laviecambot, A.; Cantuel, M.; Leydet, Y.; Jonusauskas, G.; Bassani, D.; Mcclenaghan, N. Coord. Chem. Rev. 2008, 252 (23–24), 2572–2584. Knorn, M.; Rawner, T.; Czerwieniec, R.; Reiser, O. ACS Catal. 2015, 5 (9), 5186–5193. To, W.-P.; Tong, G. S.-M.; Lu, W.; Ma, C.; Liu, J.; Chow, A. L.-F.; Che, C.-M. Angew. Chem. Int. Ed. 2012, 51 (11), 2654–2657.

1.11

Organometallic Chemistry of NHCs and Analogues

Liang Denga and Zhenbo Mob, aState Key Laboratory of Organometallic Chemistry, Shanghai Institute of Organic Chemistry, Chinese Academy of Sciences, Shanghai, PR China; bState Key Laboratory and Institute of Elemento-Organic Chemistry, College of Chemistry, Nankai University, Tianjin, PR China © 2022 Elsevier Ltd. All rights reserved.

1.11.1 Introduction 1.11.2 Overview of N-heterocyclic carbenes 1.11.3 Structures and properties of NHCs 1.11.4 Electronic properties of NHCs 1.11.4.1 Tolman electronic parameter 1.11.4.2 Lever electronic parameter 1.11.4.3 NMR chemical shift methods 1.11.5 Quantifying the steric properties of NHCs 1.11.6 Tuning the electronic and steric properties of NHCs by structure modification 1.11.6.1 NHCs with various nitrogen substituents 1.11.6.2 NHCs with diverse ring size and backbone structure 1.11.6.3 NHCs having different heterocycles 1.11.7 Chelating NHC ligands 1.11.8 Coordination compounds of NHCs 1.11.8.1 Nature of the metal-carbon(carbene) bonds 1.11.8.2 Transition-metal NHC complexes 1.11.8.3 NHCs in main group element chemistry 1.11.9 Reactions on NHC ligands 1.11.10 Organometallic chemistry of NHC analogues 1.11.10.1 Group 13 element(I) N-heterocycles 1.11.10.2 Group 14 element(II) N-heterocycles 1.11.10.3 Group 15 element(III) N-heterocycles 1.11.11 Summary Acknowledgments References

1.11.1

339 339 341 341 341 343 343 345 346 346 348 348 350 352 352 352 355 357 358 358 359 359 360 360 360

Introduction

The chemistry of carbenes has evolved in remarkable ways since the isolation of the first stable N-heterocyclic carbenes (NHCs) in 1991.1 In the past three decades, a large number of stable NHCs has been prepared. The tunable electronic and steric properties of NHCs have enabled their numerous applications in organometallic chemistry. Meanwhile, isovalent analogues of NHCs with the other p-block elements as coordinating atoms have also been prepared and used as ligands. This article gives an overview of complexes containing N-heterocyclic carbenes and their analogues. Considering the large numbers of reported NHCs and their analogues, we do not present a detailed comprehensive review on them, but instead summarize their chemistry in terms of their major categories, the synthetic methods, quantifying and tuning their electronic and steric properties, and their representative applications. A number of books and reviews on different areas of the chemistry of NHCs have also been published.2–86

1.11.2

Overview of N-heterocyclic carbenes

Carbenes form a class of neutral compounds featuring a divalent carbon atom.87 The unusual electronic properties of carbenes, which have an uncompleted octet, rendered them a great synthetic challenge to synthesize.88–97 Early attempts to synthesize the parent carbene H2C: were fruitless.98 The continuing efforts in this field led to the discovery that the stability of carbenes was strongly enhanced by heteroatom substituents.99–104 However, the unambiguous isolation of a free carbene was not achieved until the milestone report of an acyclic (phosphino)(silyl)carbene (Fig. 1) in 1988.105 The push-pull effect of the phosphino and silyl substituents plays a key role in stabilizing this carbene species.106,107 N-Heterocyclic carbenes (NHCs) are carbenes containing at least one nitrogen substituent and having the carbene center incorporated into a nitrogen heterocycle. The study of NHCs was initiated by the attempt to synthesize 1,3-diphenyl-imidazolidin2-ylidene via thermolytic elimination of chloroform from 1,3-diphenyl-2-trichloromethylimidazoline at the beginning of the 1960s.108 The thermolysis reaction gives rise to a bis(imidazoline), which provides evidence for the formation of the free NHC that, once formed, undergoes dimerization. In spite of the difficulty of accessing free NHCs, metal complexes of NHCs can be

Comprehensive Organometallic Chemistry IV

https://doi.org/10.1016/B978-0-12-820206-7.00027-5

339

340

Organometallic Chemistry of NHCs and Analogues

Fig. 1 The first isolated carbenes.

prepared by methods circumventing the use of free NHCs and had been known as early as 1960s.109–112 In 1991, the first free NHC, 1,3-bis(1-adamantyl)-imidazol-2-ylidene (IAd) (Fig. 1), was synthesized, which does not decompose even at its melting point (240 oC) under an inert atmosphere.1 The remarkable stability of IAd is attributed to both electronic and steric effects. The interaction of the lone pair electrons of the nitrogen substituent with the vacant carbene orbital reduces the electrophilic character of this special carbene. Meanwhile, the bulky N-adamantyl groups help to kinetically stabilize the carbene center with respect to dimerization. After this seminal finding, an explosion of studies on the synthesis, structure characterization and application of free NHCs was unleashed. In the area of organometallic chemistry, NHCs are welcomed as carbon-based, strongly electron-donating ligands that can form stable complexes with almost every metal in the periodic table.2–43 NHC ligands are highly modular via tuning the nitrogen substituents and the structure of the heterocyclic ring.44–51 Most of the reported NHCs are derived from imidazole (type A in Fig. 2), imidazoline (type B), benzimidazole (type C), and 1,2,4-triazole (type D), in which the carbene center has two adjacent nitrogen atoms.44 Apart from the most widely studied five-membered ring carbenes, NHCs derived from four-membered, six-membered, and even ten-membered heterocycles have also been prepared.44 An unique subset of NHCs bearing six-membered heterocycles is the N,N0 -diamidocarbenes (type F, DACs).76–79 DACs have been found to exhibit strong electrophilicity, which enables them to active small molecules such as alkenes and ammonia.113–116 NHCs with a single nitrogen substituent could also be isolated. Replacement one nitrogen atom with a sulfur atom gives thiazolin2-ylidene (type E).117 Bertrand et al. reported a series of NHCs with only one nitrogen substituent termed cyclic(alkylamino) carbenes (type G, CAAC).118–121 CAACs have recently received attention because of their even stronger electron donating and better p-accepting capability than other NHCs.118 NHCs with unique mesoionic resonance structures (type H, aNHC; type I, MIC; type J, rNHC), for which no neutral resonance form can be drawn, have been successfully isolated.122–128 Moreover, by suitable modification of the synthetic route, NHCs with chiral N-substituents (type K) or backbones (type L) can be easily prepared, which opened a route to the development of chiral NHCs for various enantioselective transformations.58–60 The most common route to free NHCs is the deprotonation of the corresponding imidazolium salts with strong bases, such as NaH, KH and KOtBu.129,130 Sterically demanding bases such as lithium diisopropylamide (LDA) and sodium bis(trimethylsilyl) amides (NaHMDS) are usually used to make CAACs118 and DACs,114 because the conjugate acid byproducts are bulky amines

Fig. 2 Representative types of NHCs.

Organometallic Chemistry of NHCs and Analogues

341

Fig. 3 Representative synthetic routes to free NHCs.

HN(iPr)2 and HN(SiMe3)2 that do not suffer NdH oxidative addition to the highly reactive carbene center. The second method was the desulfurization of imidazol-2-thiones with potassium in boiling THF.131 This method is usually applied for the preparation of free NHCs with simple aliphatic N-substituents (Me, Et and iPr). Other alternative methods to prepare NHCs have also been developed including the elimination of small molecules such as methanol,132 pentafluorobenzene,133 or carbon dioxide134 from the nascent carbene center, and the reduction of a chloroformamidinium salt with Hg(SiMe3)2135. Fig. 3 compiles these synthetic routes.

1.11.3

Structures and properties of NHCs

In parallel to the synthesis of new NHCs, spectroscopic and theoretical studies on NHCs have been performed to probe their structural and electronic properties. The X-ray photoelectron spectrum of 1,3-di-tert-butylimidazol-2-ylidene suggests the presence of both a lone pair and an empty p orbital on the carbene center.136–138 Density functional theory (DFT) studies on a series of imidazol-2-ylidenes showed that the singlet states are the ground spin states and their singlet-triplet gaps are in the range 65–85 kcal/mol.139–141 The large singlet-triplet energy gaps of NHCs are thought to originate from the inductive effect and the mesomeric effect of the nitrogen substituents. The inductive effect of the adjacent electron-withdrawing nitrogen atom significantly lowers the HOMO energy compared to the simple triplet methylene (Fig. 4), thus stabilizing the lone pair s orbital. On the other hand, the mesomeric effect allows the pp-pp delocalization of the lone pairs of nitrogen atoms into the empty p orbital of the carbene center, consequently increasing the relative energy of the LUMO.12,16 Single X-ray diffraction studies showed that the NdC(carbene) bond lengths in NHCs lie between a single and a double bond, which is in agreement with the presence of pp-pp delocalization.136,137 In addition to the electronic stabilization provided by the nitrogen atoms, the nature of the backbones (saturated or unsaturated), the ring size and the N-substituents also play important roles in stabilizing the carbene center. DFT studies on model NHCs demonstrated that imidazol-2-ylidenes show aromatic character with cyclic pp delocalization, which explains the higher stability of imidazol-2-ylidenes than imidazolidin-2-ylidenes.142,143 The NHC ring size is related to the electronic and steric properties of NHCs, which in turn affects their stability.52 Moreover, the nitrogen substituents determining the steric properties of NHCs are crucial in kinetically stabilizing NHCs by sterically preventing the dimerization process.52

1.11.4

Electronic properties of NHCs

The synthetic effort on free NHCs has led to the availability of versatile NHCs with variable steric and electronic properties. It is important to quantify the stereoelectronic properties of NHCs, which allows the judicious selection of NHC ligands for different applications. Several methods have been developed to parameterize the electronic donating and p-accepting nature of NHCs, including Tolman electronic parameter (TEP), Lever electronic parameter (LEP), NMR chemical shift methods, and DFT calculations.52–57

1.11.4.1

Tolman electronic parameter

The most commonly used descriptor to quantify the electronic properties of NHCs is the Tolman electronic parameter (TEP), which was initially applied to evaluate the electronic properties of phosphines.144,145 The basis for this quantification is the carbonyl IR

342

Organometallic Chemistry of NHCs and Analogues

(A)

(B)

Fig. 4 (A) Stabilization effects of NHCs; (B) singlet-triplet gaps of carbenes.

stretching frequencies measured for transition-metal carbonyl complexes, which is influenced by the electron richness of co-ligands. A more strongly electron donating co-ligand leads to a more electron-rich metal center, enhancing p-back donation to CO and resulting in a lower CO stretching frequency. The TEP values of various NHCs have been obtained by measuring the carbonyl A1 IR stretching frequencies of the Ni(0) complexes [Ni(CO)3(NHC)],146–158 and averaged nCO values of Ir(I) and Rh(I) complexes [MCl(CO)2(NHC)] (M ¼ Ir and Rh).159–168 It has been pointed out that care must be taken when using TEPs to compare the electronic properties of different NHCs, because the steric properties of NHCs and the measurement conditions might also influence the TEP values.52,53 The TEP values of representative NHCs measured using [Ni(CO)3(NHC)] are displayed in Fig. 5. The most commonly employed imidazol-2-ylidenes and imidazolidin-2-ylidenes (1a–1h, 2047.8–2051.1 cm−1) have lower TEP values than that of the strongest electron-donating trialkylphosphine PtBu3 (2056 cm−1), revealing that NHCs are generally stronger electron donors compared to phosphines.152 The TEP values of NHCs are inherently related to the factors that influence the electron properties of NHCs such as the number of nitrogen atoms, the ring size, the class of heterocycle, the nature of N-substituents and backbones.52,53 The replacement of the electron-withdrawing nitrogen atom by the carbon atom in types G-J increases the electron donating ability of NHCs, which decreases TEP values obtained for 1,2,4-triazolin-2-ylidenes (1i, 2057.3 cm−1), imidazol-2-ylidenes (1e, 2049.6 cm−1), and CAAC (1n, 2053.7 cm−1).169 The NHCs with expanded rings are stronger donors than the five-membered NHCs as indicated by the decrease of the TEP values in the series of five-, six-, and seven-membered NHCs (1g, 2050.8 cm−1; 1j, 2042.6 cm−1; 1k, 2041.9 cm−1).170 NHCs with alkyl backbone substituents tend to be more electron-donating (1d, 2047.8 cm−1) than the simple derivatives bearing hydrogen atoms (1c, 2050.3 cm−1).171 The presence of carbonyl groups in DAC results in the decrease of the electron density at the carbene center, which in turn increases the TEP value (1l, 2057.5 cm−1).114 Mesoionic NHCs

Fig. 5 TEP values measured with selected NHCs.

Organometallic Chemistry of NHCs and Analogues

343

are more electron-rich than the classical NHCs and possess lower TEP values (1o, 2038.4 cm−1; 1p, 2046.1 cm−1).172,173 The electronic properties of the nitrogen substituents have a strong impact on the electron donating ability of NHCs. NHCs with alkyl nitrogen substituents are generally more electron rich than the aryl substituted counterparts. The TEP values of the unsaturated NHCs (1e, 2049.6 cm−1; 1f, 2050.2 cm−1) were found to be lower than those of their saturated counterparts (1g, 2050.8 cm−1; 1h, 2051.1 cm−1).152 The same trend was also observed for the aNHCs and CAACs (1o, 2038.4 cm−1; 1n, 2053.7 cm−1).

1.11.4.2

Lever electronic parameter

The redox potential of a metal complex is inherently related to the electron-richness of the complex.174–176 For complexes without redox active ligands, the presence of a strong electron-donating ligand on a metal’s coordination sphere can increase the electron density at the metal center, therefore making the complex more easily oxidized (more negative redox potential). The most commonly used electrochemical parameter to gauge the electronic nature of ligands is the Lever electronic parameter (LEP), which evaluates the electronic properties of nitrogen, oxygen and halide donors based on the redox potentials of RuII/III couples.174–176 However, because of the difficulty of accessing the series of ruthenium NHC complexes with reversible or quasi-reversible redox potentials, the LEP parameters of NHCs were obtained from iridium and rhodium complexes [MCl(COD) (NHC)] (M ¼ Ir and Rh). The rhodium and iridium NHC complexes [MCl(COD)(NHC)] (M ¼ Ir and Rh) are relatively easy to synthesize and have reversible MI/II redox couples,164,167 which proved to be suitable redox probes for systematically comparing the electron donating abilities of NHCs.177–181 As shown in Fig. 6, the iridium complexes with imidazol-2-ylidenes and imidazolin-2-ylidenes have less positive redox potentials (in the range +0.591 V to +0.862 V) than that of complex containing a trialkylphosphine PCy3 (+0.948 V), again revealing that NHCs are stronger donors.164 The nitrogen substituents have a strong effect on the redox potentials. In the series of saturated and unsaturated NHCs bearing H, NEt2 or Br on the para-position of the N-aryl substituents, an electron-donating NEt2 group results in less positive redox potentials, and the electron-withdrawing Br group leads to more positive redox potentials. The redox potentials of iridium complexes bearing saturated imidazolin-2-ylidenes (2d, +0.759 V; 2e, +0.591 V; 2f, +0.838 V) are smaller than their unsaturated counterparts (2a, +0.786 V; 2b, +0.648 V; 2c, +0.862 V), demonstrating that imidazolin-2-ylidenes are more electron donating than imidazol-2-ylidenes.164 Significant variation in the redox potential is observed with changes to the backbone of NHC. For the six-membered NHC series, decoration of the backbone with one or two carbonyl groups leads to an increase of the redox potentials (2g, +0.61 V; 2h, +0.77 V; 2l, +0.94 V), reflecting a decreased electron-donating ability.181

1.11.4.3

NMR chemical shift methods

13

The C NMR chemical shifts of the carbene carbon atom of NHCs are diagnostic and generally appear in the downfield region between 200 and 330 ppm.182 It has been found that the singlet-triplet energy gap of NHC has a strong impact on the magnitude of the downfield shift: a smaller singlet-triplet energy gap gives a more downfield-shifted 13C NMR signal for the carbene carbon.182 Thus, the 13C NMR shifts of carbene atoms in 1,2,4-triazolin-2-ylidenes (210–220 ppm), imidazol-2-ylidenes (235–260 ppm), and CAACs (> 300 ppm) increase along with the decrease of the singlet-triplet separation. Early studies indicated that the 13C NMR

Fig. 6 Redox potentials of selected cis-[IrCl(COD)(NHC)] complexes.

344

Organometallic Chemistry of NHCs and Analogues

signal of the carbene carbon is shifted significantly upfield upon complexation and strongly influenced by the trans co-ligand.183,184 In general, a more strongly electron donating trans ligand weakens the metal-carbene bond, giving rise to a downfield-shifted 13C NMR signal. Based on this knowledge, Huynh and co-workers suggested using the 13C NMR shifts of the carbene carbon of the 1,3-diisopropylbenzimidazol-2-ylidene (1c) ligand in palladium(II) complexes trans-[PdBr2(NHC0 )(NHC)] (NHC0 ¼ 1,3-diisopropylbenzimidazol-2-ylidene) for evaluating the electronic properties of NHCs, giving a value commonly known as the Huynh electronic parameter (HEP).185,186 When the desired trans–type palladium complexes are difficult to access, the linear anionic gold NHC complexes [Au(NHC0 )(NHC)][BF4] can be used to evaluate the 13C NMR signals of 1c. The chemical shifts of the carbene carbon signals in the two type of complexes correlate according to Eq. (1).187–189 Notably, the high sensitivity of the 13C NMR based method allows the differentiation of the electron-donating ability of NHCs having similar nitrogen substituents (Fig. 7). dC ðAuÞ to HEP : HEP ¼ 1:19½AuŠ − 45:0

(1)

In addition to their strong s-donating ability, both experimental and computational studies have shown that the p-accepting strength of NHCs could be non-negligible.190–192 NMR spectroscopy is also utilized to evaluate the p-accepting ability of NHCs. The chemical shifts of 31P NMR resonances of NHC-phenylphosphinidene adducts (NHC]PPh) were initially used to gauge the p-accepting ability of NHCs.193 In a similar vein, 77Se NMR chemical shifts of NHCdselenium adducts (NHC]Se) have been used to probe the p-accepting ability of NHCs.194,195 These NHC adducts (NHC-E, E ¼ PPh or Se) have two extreme resonance structures: a resonance structure has a C ! E dative bond and a resonance structure has a C]E double bond (Fig. 8). NHCs with weak p-acceptor ability lead to an electron rich phosphine or selenium center with an upfield 31P and 77Se NMR chemical shift. On the other hand, more p-accepting NHCs result in downfield 31P and 77Se NMR signals.193–195 By measuring the dP and dSe chemical shifts of the NHC-phosphine and NHC-selenium adducts, the p-accepting properties of numerous NHCs have been quantified.193–201 The 31P NMR data correlate very well with the 77Se NMR data. The dP and dSe signals of the unsaturated NHC adducts 1e, dP ¼ −23.0 ppm, dSe ¼ 35 ppm, Fig. 9) are at higher field than those of the saturated NHC adducts (1g, dP ¼ −10.4 ppm, dSe ¼ 116 ppm), revealing that the imidazolin-2-ylidenes are better p-acceptors. CAACs are much more p-accepting than the normal NHCs as indicated by the highly downfield-shifted 31P NMR signal (1n, dP ¼ 68.9 ppm). The replacement of one nitrogen atom with one carbon atom reduces the pp delocalization of the lone pairs of nitrogen atoms into the empty orbital of the carbene center, which in turn results in better p-accepting capability of the ligand. The ring backbone also influences the p-accepting capability of NHCs. For example, introducing carbonyl groups in ligands of type F withdraws the nitrogen p-electrons into the vacant p (C]O) orbitals, and therefore reduces the p-donation to the carbene center, which in turn makes DAC a stronger p-acceptor as reflected by the highly downfield dP and dSe chemical shifts (4a, dP ¼ 78.6 ppm, dSe ¼ 856 ppm). The parameters used to describe the electronic nature of NHCs, TEP, LEP, HEP, 31P and 77Se NMR chemical shifts come from measuring the spectroscopic features of NHC-containing compounds. As the methods are differentially influenced by sigma, pi, and steric effects, the data do not always correlate well with each other. In addition to these experimental methods, DFT calculation studies on metal-NHC complexes have also been used to investigate the electronic nature of NHCs,202–210 which enables not only the systematic studies on the electron-donating nature of NHCs, but also detailed analysis on the nature of their metal-carbon bonds (vide infra).

Fig. 7 HEP values of selected NHCs.

Organometallic Chemistry of NHCs and Analogues

345

Fig. 8 The extreme resonance structures of NHC-E adducts.

Fig. 9

1.11.5

31

P and 77Se NMR chemical shifts of selected NHC adducts.

Quantifying the steric properties of NHCs

It is also important to establish a common scale to examine the steric properties of NHCs, as the steric bulk of NHCs has a profound influence on the stability and chemical reactivity of NHC-metal complexes.211,212 Though the Tolman cone angle is most often used for describing the steric nature of phosphine ligands,145 the NHCs are less cone-shaped, which has led research to develop the percentage buried volume (%Vbur) to describe the steric demanding nature of NHCs (Fig. 10a).213–216 The percentage buried volume (%Vbur) is the relative volume that the NHCs occupied in the coordination sphere of a metal center with a defined sphere of radius of 3.5 A˚ and metal-carbene bond distance of 2.00 or 2.28 A˚ , presuming free rotation of the NHC ligand along the metalcarbon(carbene) bond (Fig. 10).213–216 The %Vbur values can be calculated from X-ray crystal structure data or DFT optimized structures of free NHCs or NHC complexes using free software (SambVca).214 As the %Vbur value depends on the metal center, the ancillary ligands, and the coordination geometry, two-coordinate gold complexes [AuCl(NHC)] are often used to standardize %Vbur for NHCs because their linear geometry allows NHCs to adopt their preferred conformations.217–219 The steric bulk of various NHCs has been evaluated systematically using the percent buried volume.217–219 The %Vbur values of selected NHCs with different nitrogen substituents, NHC rings and backbone substituents in [AuCl(NHC)] are shown in Fig. 11. For the NHCs with N-alkyl substituents, the %Vbur dramatically increases when the size of alkyl groups is increased from methyl (5a, 26.3) and isopropyl (5b, 27.5) to tert-butyl (1b, 39.6) and adamantyl (1a, 39.8).213 The substitution of the NHC backbone has little influence on the steric properties of NHCs as indicated by the comparable %Vbur values of the N-2,6-diisopropylphenylsubstituted NHCs with different substitution on the backbones (1f, 45.4; 5c, 44.4; 5d, 44.9).220–222 The saturated NHC (1h, 47.0) is found to be slightly more bulky than its unsaturated NHC counterpart (1f, 45.4) as the saturated NHC ring leads to a greater NdCdN angle, which in turn places the nitrogen substituents closer to the metal center.223 Similarly, increasing the NHC ring size results in the bending of nitrogen substituents towards the metal center, thus the %Vbur values of the saturated NHCs increase (1h, 45.4; 5e, 50.9; 5f, 52.7) when expanding the ring size from five to six and seven membered cycles.214 The rigidity of the NHC ring also impacts the steric properties of NHCs. For example, the bis(oxazoline)-derived NHC (5g) is sterically demanding with a high % Vbur value (44.7).224,225 In addition, the replacement of one nitrogen substituent with a quaternary sp3 carbon substituent increases the steric hindrance, leading to a high %Vbur value for CAAC (5h, 51.0).226,227 (A)

(B)

Fig. 10 The buried volume (A), and a contour plot of a steric map of a NHC-M fragment (B). Falivene, L.; Credendino, R.; Poater, A.; Petta, A.; Serra, L.; Oliva, R.; Scarano, V.;Cavallo, L. Organometallics 2016, 35, 2286–2293.

346

Organometallic Chemistry of NHCs and Analogues

Fig. 11 %Vbur for selected NHCs in [(NHC)AuCl].

The percent buried volume is useful in evaluating the overall steric bulk of NHCs. However, it is too simple to describe asymmetric distributions of bulk around the metal center. In order to address this problem, Cavallo and co-workers have used steric maps to represent the distribution of the steric bulk of NHCs throughout the coordination sphere, wherein the steric bulk of NHCs is divided into quadrants and the percent buried volume of each quadrant is calculated (Fig. 10).215,216 The steric map is visualized as colored contours, which allows easy identification of the distribution of the steric bulk in each quadrant. For example, the steric map of a Ru-NHC complex (NHC)(Cl)2Ru]CH2 is shown in Fig. 10.216 The deep red area on the left zone shows that the mesityl group opposite the Ru]C bond shields the vacant coordination position on the Ru center. The mesityl group in the right zone of the steric map has less steric impact. Thus, the combined use of the percent buried volume and steric map can give the detailed information on the steric nature of NHC ligands, and these have become useful tools in understanding the selectivity and reactivity in catalytic reactions.228,229

1.11.6

Tuning the electronic and steric properties of NHCs by structure modification

One of the features that has led to the wide usage of NHCs in modern chemistry is the ease of modifying the structure of NHCs through the variation of nitrogen substituents, the ring backbone, the ring size, and the class of core heterocycle. This allows the preparation of NHC ligands with a wide range of electronic and steric properties.

1.11.6.1

NHCs with various nitrogen substituents

It is straightforward to tune the properties of NHCs by varying the nitrogen substituents. More electron-rich nitrogen substituents increase the electron donating ability of NHCs, and bulky nitrogen substituents exert a large steric effect on NHCs.52–57,211,212 The prototypical NHCs are the ones with hydrogen atoms as the nitrogen substituents, namely protic N-heterocyclic carbenes (pNHCs).61–64,230–232 While the free pNHCs bearing one or two NH substituents are less stable, as the proton on nitrogen often migrates to the carbene center to form more stable azole isomers, a variety of metal complexes (e.g. 6a–I in Fig. 12) bearing pNHC ligands have been prepared and extensively studied.233–250 The metal complexes featuring pNHCs are special, because the reactive NH group could form hydrogen bonds with substrates, which can aid in metal-ligand cooperative bond activation.64 The acidic proton on nitrogen can be deprotonated to form azolyl complexes, which provides potential for N-functionalization of the pNHC complexes and also metal-ligand cooperative bond activation.250–259 Varying the nitrogen substituents has enabled the preparation of a vast number of NHCs (Fig. 13). The nature of the nitrogen substituents has significant influences on catalyst activity of their metal complexes. The properties of NHCs with alkyl N-substituents are usually modified by changing the number of substituents on the N-bound carbon. NHCs with alkyl N-substituents have been used for the stabilization of cobalt(I) complexes, and homoleptic complexes of the type [Co(NHC)n]+ (n ¼ 4, 3, 2) proved accessible as the NHC ligand changes from 1,3-dialkyl-4,5-dimethylimidazol-2-ylidenes (IR2Me2, 7a) to 1,3-dicyclohexylimidazol-2-ylidene (ICy, 7b) and to IAd, which parallels with the increased steric demanding nature of the alkyl N-substituents.260–263 The most common approach to modify NHCs bearing aryl N-substituents is to vary the para- and ortho-substituents of the aryl groups. The ortho substituents on the aryl group of an NHC exert a strong steric effect. For example, the replacement of the methyl groups of IPr by phenyls leads to a highly hindered IPr ligand (7c, IPr ¼ 1,3-bis(2,6-bis (diphenylmethyl)-4-methylphenyl)imidazo-2-ylidene), which has been shown to be useful in a number of transition-metal catalyzed reactions.264–266 The para substituents on the aryl group of an NHC also have non-negligible steric influence. A prominent

Organometallic Chemistry of NHCs and Analogues

347

Fig. 12 Selected examples of metal pNHC complexes.

Fig. 13 Selected examples of modulation of N-substituents of NHCs.

example is the IPr derivatives with large groups at the para position (7d–f). Palladium complexes with these NHCs are more active precatalysts for Suzuki-Miyaura coupling than that with the parent IPr.267 The modulation of the para substituent on the phenyl ring of NHCs can be used for electronic effects: for example, the Grubbs-type ruthenium catalysts bearing NHCs with different para groups (7g–j) show a significant influence on the redox potential.268 Unsymmetrical NHCs with different nitrogen substituents were also prepared in order to tune their electron-donating ability and the steric bulk. The ruthenium complex bearing the unsymmetrical NHC ligand 7k catalyzes highly Z-selective cross-metathesis reactions.269 The cross-metathesis reactions having the ruthenium complexes bearing N-aryl-N0 -benzyl substituted unsymmetrical NHC ligands 7l shows slower initiation but is more selective than conventional Grubbs-type ruthenium catalysts.270 Nickel bearing the N-aryl-N0 -cyclohexyl NHC 7m is useful for regio- and diastereoselective cross-hydroalkenylation of endocyclic dienes with -olefins.271 A significant part of research in modification of nitrogen substituents of NHCs has been directed toward the development of chiral NHCs.58–60 While introducing chirality on the NHC backbone is a useful way of constructing chiral NHC ligands, NHCs with chiral nitrogen substituents could have their chiral groups being closer to the metal center, and hence are expected to be more effective in asymmetric induction. The availability of numerous primary amines containing chiral substituents lays the foundation for the structural diversity of this type chiral NHCs. As an early exploration on this aspect, the rhodium complexes with chiral NHC ligands 8a and 8b were found to catalyze the hydrosilylation of ketones with moderate enantioselectivities.272–274 Further exploration in the field then led to the recognition that the performance of the chiral NHCs in enantioselective transition-metal catalysis is related to the size and rigidity of the nitrogen substituents. Representative chiral NHCs of the type are shown in Fig. 14, including the ones bearing bulky alkyl substituents 8c275 and 8d276, rigid cyclophane moieties 8e,277 bulky phenyl groups having chiral benzylic groups at the 2,6-positions 8f,278,279 binap group 8g,280 as well as fused rings 8h,281 8i,282 8j,283 8k284 and 5g.225,285

348

Organometallic Chemistry of NHCs and Analogues

Fig. 14 Selected NHCs with chiral nitrogen substituents.

1.11.6.2

NHCs with diverse ring size and backbone structure

Great efforts have also been exercised in the development of NHCs with ring sizes beyond the classical five membered ring and also ones featuring functionalized backbones. NHC ligands with core rings varying from four- to ten-membered rings, e.g. 9a,286 9b,287,288 9c,289–292 9d,293,294 9e295 and 9f,295 are known (Fig. 15). NHCs with expanded ring size generally have larger NdC (carbene)dN angles as compared to those in the five- and six-membered NHCs, and the C(R)dN and NdC(R) bonds are not coplanar. Therefore, expanding the ring size generally increases the steric bulk of NHCs. The substitution of the hydrogens at the backbone of IMes (1,3-dimesityl-imidazol-2-ylidene) by chlorides gave 1,3-dimesityl4,5-dichloroimidazol-2-ylidene (9g), which is air stable.296 The incorporation of photochromic dithienylethenes into the backbone of IMes resulted in a photo-switchable NHC through reversible ring-opening and ring-closing, which was used to reversibly tune the electronic properties of the carbene center.297–300 The p-electron-withdrawing carbonyl groups on diamidocarbenes (DACs) enhance the electrophilicity for DACs and result in reactivity like triplet carbenes.301 Recently, anionic NHCs (9j) with a weakly coordinating anionic borate moiety in the backbone have been reported.302,303 The synthesis of imidazolin-2-ylidenes can use chiral 1,2-diamines as the starting material, which produces NHCs with chiral backbones. Among this type of chiral NHC ligands, the diphenyl-substituted imidazolin ylide ligands are most well-developed. As the chiral backbones are far from the carbene carbon centers, bulky nitrogen substituents must be incorporated in this type of chiral NHC ligands (e.g. 9k–9m) to achieve high asymmetric induction in metal-NHC-catalyzed reactions. Fused ring systems generally are more rigid,304–308 and this strategy has been used toward chiral bicyclic and tricyclic NHCs, e.g. 9n,309 and 9o.310 The placement of one nitrogen atom in the bicyclic backbone restricts the delocalization of its lone pair to the carbene center in 9p, thus increasing the electrophilicity of its NHC.311 As another interesting type of NHCs with special backbones, ferrocene-based NHC ligands 9q can be viewed as a type of NHCs having six-membered core rings.312–314 When coordinating with transition-metals, long-range electronic communication between the iron center, carbene carbon, and transition-metal was envisioned.312

1.11.6.3

NHCs having different heterocycles

While 1,3-imidazol-2-ylidene and 1,3-imidazolin-2-ylidene are the most well-studied NHC ligands, a large number of other NHC ligands have also been developed, which show different electronic and steric properties (Fig. 16). Variations of 1,3-imidazol2-ylidenes with one NR group replaced with an oxygen or sulfur atom are oxazolylidenes (10a) and thiazolylidenes (E), which are

Organometallic Chemistry of NHCs and Analogues

349

Fig. 15 Representative NHCs with different ring size and backbone.

Fig. 16 NHCs with different heterocycles beyond imidazol-2-ylidenes.

less sterically demanding.315–322 Free oxazolylidenes are generally unstable, but their metal complexes are known. Thiazolylidenes are expected to have strong p-accepting capability as the smaller mesomeric effect of sulfur atom leads to reduced pp-pp delocalization.44 NHCs featuring three nitrogen atoms can display very different electronic properties that can be influenced by the location of the nitrogen atoms. In comparison to imidizolin-2-ylidene, 1,2,4-triazolin-5-ylidenes (D) are weaker donors due to the inductive effect of the third nitrogen.44 However, 1,2,3-triazolin-5-ylidenes with one nitrogen adjacent to the carbene center (I) have stronger electron donating ability.122 Similarly, “abnormal NHCs” with the carbene center located at C4 or C5 of the imidazolium ring, J and H, respectively, are also particularly strong donors because they have a carbon rather than a nitrogen adjacent to the coordinating carbon (see Section 1.11.4.1 above). The isolation of such NHCs are more challenging, but their application as ligands in organometallic chemistry is flourishing.51

350

Organometallic Chemistry of NHCs and Analogues

One of the most significant advances in the design of NHCs has been the development of cyclic (alkyl)(amino)carbenes (CAACs, G).19 The replacement of one electron-withdrawing nitrogen substituent of NHCs by an electron-donating alkyl group in CAACs reduces the inductive effect and also lowers the degree of pp delocalization, and therefore CAACs are more nucleophilic and electrophilic than the classical NHCs.19 The properties of CAACs can be further tuned by changing the ring size and the ring backbone similar to that observed in classical NHCs.323,324 The six-membered CAACs 10b and 10f provide more steric hindrance than G. They are also expected to have enhanced donor and acceptor properties.323,324 Computational studies suggest that the cyclic (alkyl)(amido)carbenes 10e325 and cyclic (amino)(aryl)carbenes 10c326 and 10d327 are stronger p-acceptors as compared to G. So far, free CAACs of G and 10b have proven accessible, and the CAACs 10c–10f are only known as ligands in metal complexes.19,328,329

1.11.7

Chelating NHC ligands

There are many chelating ligands featuring NHC units.65–71 They enjoy the strongly electron-donating nature of NHCs and also the chelating effect that further improves the stability of metal NHC complexes. Bis-NHCs,330–338 tris-NHCs,339–345 and tetra-NHCs346–352 present a special type of chelating NHC ligands (Fig. 17). They can be prepared by substitution of imidazoles with halide-substituted linkers or by ring-closure reactions of the pre-assembled amino units with pre-carbenic units.65–71 While the isolation of free chelating NHC ligands is often challenging, their metal complexes are easily accessible. The methylene bridged bis-NHC ligands 11a, the chiral BINAP bridged bis-NHC ligand 11b and the borate bridged NHC ligand 11c are useful ligands in transition-metal catalysis.353–360 Tripodal tris-NHCs 11d–11f are analogues of tris(phosphine) ligands and have shown wide application in stabilizing metal complexes in unusually high oxidation states and featuring metal-ligand multiple bonds.339–345 The macrocyclic tetra-NHCs 11g–11i have received great attention for their structural similarity to porphyrins but possessing

Fig. 17 Selected examples of bi- and poly-dentate NHCs.

Organometallic Chemistry of NHCs and Analogues

351

stronger ligand fields.361 Late transition-metal complexes of macrocyclic tetra-NHC complexes have been subjected to extensive study,351,362–367 and among them the ability of tetra-NHC ligands in supportingiron(IV) terminal oxo,363 iron(IV) terminal imido,364 and iron alkylidene species351 is noteworthy. In addition to the utility as ligands for mononuclear metal complexes, many bis-, tris-, and tetra-NHCs, e.g. 11j–11l, have been developed for the construction of organometallic rectangles, cages, and polymers.368–396 Hybrid chelating NHC ligands, which have both NHC in addition to a different kind of donor group, are the most well-developed chelating NHC ligands.5,65–77 The combination of the advantages of NHCs and other donor groups offer vast opportunities for tuning the electronic and steric ligand environment. Hybrid chelating NHCs can be approached via several routes: substitution reactions of donor-tethered electrophiles with imidazoles, ring-closure reactions between donor-tethered amines and formal aldehyde or ortho-formate, template-controlled synthesis, and post-functionalization.5 The bi-, tri-, tetra-, penta-, and even hexa-dentate hybrid chelating ligands are all known. Fig. 18 exemplifies some hybrid chelating ligands 12a–12r.397–419 The additional donor groups can be carbanion, alkene, amine, imine, pyridine, amido, ether, alkoxido, carboxylate, silyl, silylene,

Fig. 18 Examples of hybrid type chelating NHC ligands.

352

Organometallic Chemistry of NHCs and Analogues

thioether, phosphine, phosphinido, and thiolate.5,65–77 A number of NHC-containing pincer ligands have also been developed, either having the NHC unit in the center,5,420–424 or with two NHC units on the sidearms.402,412–414,425–446 NHC-containing pincers are a class of strongly electron-donating ligands with high scaffold rigidity, and are widely applied in transition-metal-catalyzed reactions.447–452 In addition to the chelating ligands of imidazol-2-ylidenes, hybrid chelating NHC ligands containing CAACs, aNHCs, and MICs (e.g. 12s–u) have recently been developed.50,227,453–463 A number of reviews5,65–77 and monographs3–8 on hybrid chelating NHCs exist in literature.

1.11.8

Coordination compounds of NHCs

With their unique steric and electronic features, NHCs have become indispensable ligands in coordination chemistry. They are widely used for the preparation of transition-metal and main group element compounds. These NHC-coordinated species are widely used in diverse research areas such as catalysis, materials science, and medicinal chemistry.2–43

1.11.8.1

Nature of the metal-carbon(carbene) bonds

The nature of the metal-carbon(carbene) bond in NHC complexes has been addressed through extensive theoretical studies.464–477 Two reviews published in 2007 and 2009 have summarized the understanding.55,478 The composition of the frontier molecular orbitals of NHCs indicates that they have the potential to function as s-donating ligands that are either p-donating or p-accepting (Fig. 19).468–472 As the metal-carbon(NHC) bonds are generally much longer than metal-carbon double bonds and closer to metal-carbon single bonds, one can envision that the p-interaction between metal and NHC is not strong. Accordingly, NHCs were initially considered simply as s-donors. Later, more studies suggest that in some cases the p-interaction between metal and NHC ligand is not negligible. For example, DFT calculations suggest that the p-backdonation from the filled d orbitals of metal center into the empty pp orbital of the carbene center is a significant contribution in the NHC complexes of nd10 metals, e.g. silver(I),344 copper(I),344 gold(I),344 palladium(0),344 and nickel(0).479 In contrast, p-donation from NHC to the metal was found in the cyclometallated iridium(I)- and rhodium(I)-IBu0 complexes [MCl(IBut0 )2 and [M(IBut0 )2]+ (M ¼ Rh, Ir)].470 Compared to the metal complexes of regular NHCs, the relatively low energy of LUMO of CAAC versus that of regular NHCs leads to enhanced metalto-CAAC p-backdonation, which can be discerned from the relatively short metal-carbon(CAAC) bonds of metal CAAC complexes. As a result of this backdonation, open-shell metal-CAAC complexes might have significant spin density on the C(carbene)-N moiety.480–490 Energy decomposition analysis has been used to quantify the composition of interactions of metal-carbon(NHC) bonds.493–495 Among the attractive interactions, electrostatic attraction is found to be the dominating contributor in stabilizing the metalcarbon(NHC) bond. Calculations on the group 11 metal NHC complexes (NHC)MX (M ¼ Cu, Ag, Au; X ¼ F, Cl, Br, I) revealed that the electrostatic attraction accounts for more than 65% of the bonding interactions between NHC and MCl, whereas the contributions of the orbital interaction are less than 35%.468 For the orbital interactions, p-backdonation accounts for about 20% of the total orbital interaction energy in these coinage metal complexes.468 A systematic study on NHC complexes of early to late transition metals revealed trends in the relative amounts of s- and p-contributions toward metal-carbon(NHC) bonds of different metals.472 The s-contributions to the metal-NHC bonds decrease from 90% for d0 (Ti, Zr and Hf ) systems to 80% for d10 (Pd and Pt) systems, and the p-contributions to the orbital interaction increase from 10% to 20%, while the percentage of p-donation to the p-contribution decreases from 35 to 10%. In addition to these orbital considerations, favorable London dispersion interactions could play an important role in stabilizing metal-NHC complexes.496

1.11.8.2

Transition-metal NHC complexes

A number of methods are available for the preparation of NHC complexes.2–8 The most commonly used method to treat free NHCs, isolated or in situ generated from the corresponding azolium salt, with suitable metal precursors like their phosphine analogues. A prominent example is the synthesis of the Ru-NHC complexes trans-[RuCl2(CHPh)(PCy3)(NHC)] from the reactions of free NHCs with the first-generation olefin metathesis catalyst trans-[RuCl2(CHPh)(PCy3)2].497–499 Another widely used method is the transmetallation of NHC complexes of silver and copper with other metal compounds.397,500 Other alternative methods include

Fig. 19 Possible orbital interaction between NHC and transition metal center.

Organometallic Chemistry of NHCs and Analogues

353

Fig. 20 Selected examples of NHC-stabilized transition metal complexes.

the reactions of electron-rich enetetramines with metal species,501,502 oxidative addition of an azolium salt,503–507 a-elimination of small molecule from an azolium salt, 508–511 protonation or alkylation of azolyl complexes,512,513 and template-controlled synthesis using isocyanide complexes.514–518 NHC complexes of all transition-metals and most of the stable f-block elements have been reported, and are summarized in earlier reviews on the chemistry of metal-NHC complexes.2–43 While NHCs in many of the complexes can be viewed as surrogates of phosphines, there are ample metal complexes that take advantage of the unique steric and electronic nature of NHCs, where analogous phosphine complexes are unknown. One category of such complexes are the 3d metal-ligand multiple bond complexes (Fig. 20), e.g. the iron(V) nitrido complexes 13a,519,520 the manganese(V) nitrido complex 13b,521 the iron(IV) oxo complex supported by a tetra(NHC) ligand 13c,363 the cobalt(III) oxo complex 13d,488 the iron(IV), (V) and (VI) bismido imido complexes 13e–g,489,522,523 cobalt(IV), and cobalt(V) bisimido complexes 13h–i,524 and the two-coordinate nickel,525 cobalt526,527 and iron528 imido complexes 13j–l. Another representative category of NHC complexes are the low-coordinate low-valent metal complexes that include the two-coordinate zero-valent and monovalent metal complexes [M(NHC)2] (M ¼ Ni (13m)529, Pd530, Pt531), [M(NHC)2]+ (M ¼ Ni529, Co(13n)532, Fe533–535), [M(cAAC)2] (M ¼ Zn,487 Cu,483,536 Ni,537 Co(13o),538 Fe,484 Mn,486 Pd,539 Pt539), and [M(cAAC)2]+ (M ¼ Co(13p),538 Fe484,534), the three-coordinate zerovalent metal complexes [M(NHC) (Z2-alkene)2] (M ¼ Ni,540 Pd,541 Pt,542 Co,488,543 Rh,544 Fe,545,546 Mn(13q)491) and [M(NHC)2(Z2-alkene)] (M ¼ Ni(13r),547 Co543), as well as the three-coordinate subvalent metal complexes [M(NHC)(Z2-alkene)2]− (M ¼ Co(13s), Rh)544 and [Fe(cAAC)2(N2)]− (13t).484 NHC ligands are also used for stabilizing biologically relevant complexes, such as the mixed-valent, Fe(II)Fe(I) complex [(m-pdt)[Fe(CO)2(PMe3)][Fe(CO)2(IMes)]+(13u),548 the trinitrosyliron complex [Fe(NO)3(IMes)]+(13v)549, the all-ferrous iron-sulfur cluster [Fe4S4(IPr2Me2)4] (13w),550 and the organometallic iron-sulfur cluster [Fe4S4(IMes)3(CH2Ph)] (13x)551. A large number of metal-NHC complexes have been described in the literature. However, only a handful of them contain bridging NHC ligands with carbene carbons as bridging atoms. These include some complexes of alkali and coinage metals.542–554 In addition, a dinuclear nickel complex 13z bridged through the carbon atom of a tetradentate bis(N-imidazolylpyridine)methane ligand has been reported.555 A number of reviews on the diverse properties and catalytic applications of metal NHC complexes have been published.8–43 The application of NHC complexes in homogeneous catalysis has developed vigorously since the seminal report of a Pd-catalyzed Mizoroki-Heck reaction by Herrmann and co-workers.556,557 Particularly active sub-areas include olefin metathesis,24,27 CdC cross-coupling reactions,2,3,558 hydrogenation,2,3 reductive coupling reactions,559,560 and alkyne transformations.2,3 Fig. 21 lists examples of olefin metathesis catalysts 14a–d,268,561,562 CdC and CdN cross-coupling catalysts 14e–h,563–573 hydrogenation catalysts 14i–l,574–592 hydro-functionalization catalysts 14m–p,592–596 and catalysts for alkyne transformation 14q–r.593–612

354

Organometallic Chemistry of NHCs and Analogues

Fig. 21 Selected examples of metal-NHC catalysts.

In addition to their applications in organic synthesis, NHC metal complexes are also popular catalysts in small molecule activation reactions.613–615 For example, the iridium complex [Ir(bis-NHC)(OAc)I2] (14s) is particularly efficient for the reduction of CO2 to formate.616 The redox-active NHC-pyridine macrocycle 14t617 shows high selectivity in for electrocatalytic CO2 reduction to CO in protic solution. A dinuclear RudPd complex 14u proved effective in catalyzing photochemical proton reduction.618 Iridium NHC complexes 14v619 and 14w620 have been applied as catalysts for water oxidation. In N2 reduction reactions, a molybdenum complex containing a tridentate phosphine-NHC-phosphine ligand 14X displays high catalytic activity for reduction of N2 into NH3 with alcohol or water as solvent and SmI2 as reductant.621 Metal NHC complexes have attracted great interest in medicinal chemistry as well.622 A number of late transition-metal-NHC complexes have been found to show antimicrobial and antitumor activity. 79–82,623–628 Among them, the complexes of silver, gold, and platinum were subjected to the most extensive studies. Investigations on the biological activity of the NHC complexes of ruthenium, rhodium, copper, and nickel are relatively scarce (Fig. 22). A factor that contributes to the applicability of the NHC metal complexes in these medicinal studies is the stability in aqueous solutions. In terms of NHC ligands, while imidazol-2ylidene-based NHCs are most popular, imidazol-4-ylidene, 1,2,4-triazole-derived chiral NHC, and pyrazolin-3-ylidene, e.g.

Organometallic Chemistry of NHCs and Analogues

355

15a–c629–631 have also been employed. As the anti-cancer cell activity of a medicinal molecule require has the “Janus” requirement of water solubility and lipophilicity, NHC ligands having both polar functionalities, e.g. hydroxyl/ether, and aryl groups, e.g. 15d–e, 632,633 have been chosen judiciously in the design of medicinally oriented metal NHC complexes. Another meaningful molecular design is the incorporation of biologically active moieties in the NHC ligands as demonstrated by the studies on caffeine-derived NHC silver complex 15f,634 clotrimazole-derived NHC silver complex 15g,635 and the peptide-functionalized NHC gold complex 15h.636 Recently, the cytotoxicity of transition-metal complexes bearing NHC ligation, e.g. 15i637 and 15j,638 has also gained attention. NHCs are strong s-donating ligands with either p-donating or p-accepting ability. These electronic features can potentially render the d–d ligand-field excited states and the ligand-to-metal change-transfer states of their metal complexes high in energy and the metal-to-ligand charge-transfer states low in energy, which can endow metal NHC complexes unique photophysical and photo-chemical properties. A number of reviews on photofunctional complexes of NHCs are known.639–643 Research efforts in luminescent NHC metal complexes are mainly aimed at the development of blue color organic light emitting diode material and of phosphorescent materials for photochemical reactions, dye-sensitized solar cells, and medicinal chemistry. The explorations initially were mainly focused on the d6, d8, and d10 metal complexes, e.g. the rhenium(I) complex 16a,644 ruthenium(II) complexes 16b–d,645–647 iridium(III) complexes 16e–h,648–651 platinum(II) complexes 16i–k, 652–654 gold(I) complexes 16l–n,655–658 the silver complexes 16o–p,659,660 and the copper(I) complexes 16q–s.554,661–663 More recently, photoluminescent d5 metal complexes, e.g. the iron(III) NHC complexes 16t–u, 664,665 have been found to have a ligand-to-metal charge-transfer states with an impressively long lifetime (Fig. 23). The diverse topologies of poly-NHC ligands are beneficial for the formation of diverse metal-NHC architectures. A series of three-dimensional organometallic supramolecular assemblies featuring NHCs has been reported (Fig. 24),1,3,666–682 e.g. the trinuclear silver(I) carbene organometallic supramolecular structure 17a376, the Ag8L4-type supramolecular assemblies 17b–17e392 and the three-dimensional interlocked AgIdNHC architecture 17f396. Very recently, two hexanuclear AgIdNHC assemblies 17g–h featuring a tris-NHC precursor with two types of NHC donors were synthesized.668 Heteroligand assemblies with more than one type of donor have also been studied,84,85 e.g. an triangular hexanuclear nickel organometallic cylinder 17i386 composed by three di-NHC donors and two N-containing ligands (TPT) and a tubular octanuclear gold(I) NHC complex 17j.667 The development of NHC-functionalized heterogeneous metallic materials672–678 and MOF materials679–682 has also attracted considerable interest as NHCs could offer high stability and new prospects for nanoparticles as well as enable size/shape selectivity and recyclability of MOFs.

1.11.8.3

NHCs in main group element chemistry

The chemistry of main group compounds stabilized by NHCs have been described in a number of excellent reviews.86,683–685 The bonding interaction between a main group element and NHC is generally described as a dative covalent bond.686 The use of NHCs as ligands has allowed the preparation of a great number of unprecedented main group element species, especially those with multiply bonded and low valent main group elements (e.g. 18a–o, Fig. 25).86 Prominent examples include the diatomic species of the form NHC ! E02 NHC (E ¼ B, Si, Ge, Sn, P and As, 18a–f),687–690 monoatomic zerovalent main group element compounds (NHC)nE0 (E ¼ C, Si, and Ge, 18g–k),691–696 the first stable nucleophilic neutral boron species (CAAC)2BH (18l),697 and the first neutral zero-valent s-block complex (CAAC)2Be0 (18m) by making use of Be ! CAAC p-back bonding.492 Very recently, CAAC-supported borylene species were found to show intriguing N2 binding, reduction and coupling (18n–o).698–700 The use of CAACs for the stabilization of main-group element radical complexes of carbon, antimony, borane, aluminum, silicon and phosphorus is also extensive.685

Fig. 22 Selected examples of metal-NHC complexes for medicinal application.

356

Organometallic Chemistry of NHCs and Analogues

Fig. 23 Selected examples of photolumincent metal-NHC complexes.

R = Me

Fig. 24 Representative examples of organometallic supramolecular assemblies based on poly-NHC ligands.

Organometallic Chemistry of NHCs and Analogues

357

Fig. 25 Selected examples of NHC-stabilized low valent main group element species.

1.11.9

Reactions on NHC ligands

Though NHCs are generally inert and present as spectator ligands in compounds, examples of reactions of NHC ligands on the coordination sphere of transition-metal and main-group elements are not rare. The representative reactions of NHC ligands are reductive elimination reactions to form imidazole salts, migratory insertion reactions to yield diaminoalkyl, iminato, and urea complexes, cyclometallation reactions on the N-substituents, dealkylation and dearylation reactions of N-substituents, and ring-expansion and carbon-extrusion reactions (Fig. 26).701–706 CdC, CdH, and CdX (X ¼ hetero-atom) bond-forming reductive elimination reactions of NHC ligands were mainly observed on late transition-metal complexes. This type of reaction represents one of the catalyst deactivation pathways in metal-NHC-catalyzed cross-coupling reactions.704–710 On the other hand, it is thought to be responsible for the formation of nano-sized metal clusters or molecular catalysts derived from metal-NHC complexes that are the genuine active catalysts for reactions.711–716 Cyclometallation reactions NHC ligands are commonly observed for both early and

Fig. 26 Reactions on NHC ligands.

358

Organometallic Chemistry of NHCs and Analogues

late transition-metal complexes. The type of reaction can be viewed as a way of generating carbon-donor-functionalized NHC ligands. The latter can be further used for the preparation of other donor-functionalized NHC metal complexes.75,717–728 While most of the cyclometallation reactions of NHC ligands are achieved by CdH bond cleavage, there are also examples of CdC bond of NHCs to afford cyclometallated NHC complexes.729,730 Another interesting set of reactions on NHC ligands is the dealkylation and dearylation of N-substituents of NHCs via transition-metal-mediated CdN bond activation. This type of reaction gives C-imidazole complexes. Most of the CdN bond cleavage reactions are thought to be induced by an initial cyclometallation reaction.544,731–737 However, a recent study suggests that low coordinate cobalt(-I) and rhodium(-I) species can activate the CdN bond of N-aryl NHCs through an oxidative addition mechanism.544 Ring-expansion reactions of NHCs are mostly observed in the reactions of NHCs with silicon, beryllium and boron hydride compounds.738–742 Some carbon-extrusion reactions of NHCs are thought to come from ring-expansion intermediates,743–745 and others are thought to involve migratory insertion.746–749

1.11.10

Organometallic chemistry of NHC analogues

Along with the exploration on NHC ligands, analogous ligands having a group 13 element (B, Al, Ga, In), a group 14 element (Si, Ge, Sn), or a group 15 element (N, P, As, Sb) in place of the carbene carbon in the N-heterocycles have received increasing attention in organometallic chemistry.

1.11.10.1 Group 13 element(I) N-heterocycles Two types of isoelectronic analogues of NHCs have been developed: anionic group 13 element(I) N-heterocycles and neutral group 13 element(I) N-heterocycles.750 The first anionic borane(I) N-heterocycle was isolated as a lithium salt (19a), through reduction of the 1,3,2-diazaborole with lithium metal.751 Following this initial report, different N-heterocyclic boryl ligands (19b–d) have been prepared as lithium salts,752–754 and a variety of transition-metal boryl complexes were prepared by their reactions with transition-metal precursors.755–760 N-heterocyclic boryl ligands have also been used to stabilize unprecedented main group element species such as acyclic silylene, group 13 element(II) radicals and digermavinylidenes.761–763 The neutral N-heterocyclic aluminum(I) ligands 19e–f have been prepared and shown diverse reactivities.764–770 The neutral gallium(I), indium(I) and thallium(I) heterocycles (19i–k, 19l–n) incorporating b-diketiminate and guanidinate ligands are also known.771–774 The study of N-heterocyclic aluminyl anions is undergoing a recent renaissance. Aluminyl anions incorporating bulky amido and alkyl ligands (e.g. 19g and 19h) were shown to behave as aluminum nucleophiles for the activation of inert small molecules.775–780 The first anionic group 13 element(I) N-heterocycle was the gallyl anion 19o.781 Variants of 19o with different nitrogen substituents and backbone are known.782–784 These gallyl anions have been used as ligands to prepare transition-metal and p-block element compounds.750 So far, only one example of indyl anion 19p was reported (Fig. 27).785

Fig. 27 Selected examples of group 13 element(I) N-heterocycles.

Organometallic Chemistry of NHCs and Analogues

359

1.11.10.2 Group 14 element(II) N-heterocycles To date, a significant number of heavier group 14 analogues of NHCs are known, which are still under rapid development and have proven to be useful ligands in organometallic chemistry. 750,786 The first N-heterocyclic silylene (NHSi) 20a was reported in 1994,787 followed by the diaryl substituted version 20b, saturated version 20c, benzo-fused version 20d, six-membered version 20e, and alkyl(amino) version 20f (Fig. 28).788–794 Amidinate-stabilized silylenes 20g are the other important type of NHSis, and their incorporation into chelating ligand scaffolds has been thoroughly studied.795–804 DFT calculations investigated the electronic properties of NHSis, indicating that the s-donor and p-acceptor strengths of NHSis could exceed NHCs or phosphine ligands.805,806 The first NHSi complex implicated in homogeneous catalysis was a dinuclear NHSi Pd complex.807 Transition-metal complexes bearing chelating silylene ligands have shown excellent catalytic performance in many transformations such as Heck or Suzuki cross-coupling, alkyne cyclotrimerization, hydroformylation, CdH borylation, hydrosilylation, and hydrogenation.796–803 A large number of N-heterocyclic germylenes, stannylenes and plumbylenes have also been reported.750 However, their corresponding transition-metal complexes are rare in comparison to the plethora of NHSi complexes, which might be due to their lesser s-donating ability.

1.11.10.3 Group 15 element(III) N-heterocycles The isovalent group 15 analogues of NHCs, namely N-heterocyclic nitrenium, phosphenium, arsenium, and stibenium, bear a formal positive charge, leading to strong Lewis acidities at the central group 15 atoms.808–811 Theoretical studies on the electronic properties of their metal complexes suggest that the MdE (E ¼ N, P, As and Sb) bonding is dominated by p-back donation from the metal to ligand with weak s-donation from ligand to the metal (Fig. 29).812–814 The application of cationic species (e.g. 21a–d) as ligands for transition-metals has been developed, among which N-heterocyclic phosphenium are the most widely studied in the field of transition-metal chemistry.815–824 In most of the N-heterocyclic phosphenium complexes, the PdM s-bonding and M-to-P p-back bonding contributions result in a short PdM distance and a planar geometry at the phosphorus center. In some cases, P-to-M s-bonding contribution is rather weak and the N-heterocyclic phosphenium ligand can be viewed as Z-type ligands,

Fig. 28 Selected examples of group 14 element(II) N-heterocycles.

(A)

(B)

Fig. 29 (A) Schematic representation of bonding group 15 analogues of NHCs; (B) represent examples of group 15 elements N-heterocycles.

360

Organometallic Chemistry of NHCs and Analogues

generating a pyramidal geometry for the P atom.825 Recent studies on metal complexes with chelating N-heterocyclic phosphenium pincer ligands demonstrated that certain N-heterocyclic phosphenium metal complexes should be better viewed as transition-metal phosphido complexes.826,827 The binding ability of N-heterocyclic nitreniums is very weak due to their weak p-accepting and s-donating ability. By introducing N-heterocyclic nitrenium into a tridentate framework with two phosphine donors (21a and 21b), the coordination of N-heterocyclic nitreniums to various transition metals proved possible.828

1.11.11

Summary

NHCs have become ubiquitous and indispensable ligands in the field of organometallic chemistry. The strongly s-donating nature of NHCs, their demanding steric profile, and the ease of synthesis that allows facile modification of the electronic and steric properties of NHCs, form the basis for their wide applications. Over the years, a variety of NHCs having different heterocycles, nitrogen substituents, and ring backbones, as well as donor-functionalized NHCs, have been developed and their electronic and steric comparison to other ligand sets have been parameterized by spectroscopic and computational methods. Numerous metal and main-group element complexes with NHC ligation have been synthesized, and some of them have structural and electronic features that are hardly observed in other complexes. In organometallic catalysis, NHC metal complexes have been applied successfully as catalysts in many types of organic transformations, in which phosphine metal complex-based catalysts had previously been dominant. In other examples, reactions that cannot be done with phosphine-metal catalysts were enabled by NHC metal complexes. Finally, NHCs have now found application in modern materials science and medicinal chemistry. Again, these applications, e.g. as ligands for stabilizing gold clusters, modifying metal surfaces, constructing organometallic frameworks and rectangles, new luminescent molecules, as well as anti-tumor reagents, benefit from the unique electronic and steric property of NHC ligands. With these intriguing areas of progress, we can foresee that the continuing exploration on the metal and main-group element compounds bearing NHC ligands or their analogues should lead to continued new applications in synthetic chemistry, materials science, and medicinal chemistry.

Acknowledgments This work was supported in part by the National Natural Science Foundation of China (Nos. 21725104, 21690062, 21821002, and 22071124) and the Program of Shanghai Academic Research Leader (No. 19XD1424800). Z.M. gratefully acknowledge the College of Chemistry of Nankai University for generous financial support. We thank Prof. Ying-feng Han for the helpful discussion.

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. 29.

Arduengo, A. J., III; Harlow, R. L.; Kline, M. J. Am. Chem. Soc. 1991, 113, 361–363. Nolan, S. P., Ed.; In N-Heterocyclic Carbenes in Synthesis; Wiley-VCH: Weinheim: Germany, 2006. Glorius, F., Ed.; In N-Heterocyclic Carbenes in Transition-metal Catalysis; Springer: Berlin, Germany, 2007. Diez-Gonzalez, S., Jahnke, M., Hahn, E., Eds.; In N-Heterocyclic Carbenes: From Laboratory Curiosities to Efficient Synthetic Tools; Royal Society of Chemistry: Cambridge, UK, 2010. Kuhl, O., Ed.; In Functionalised N-Heterocyclic Carbene Complexes; Wiley-VCH: Chichester, 2010. Cazin, C. S. J., Ed.; In Heterocyclic Carbenes in Transition-metal Catalysis and Organocatalysis; Springer: Berlin, Germany, 2011. Nolan, S. P., Ed.; In N-Heterocyclic Carbenes Effective Tools for Organometallic Synthesis; Wiley-VCH: Weinheim, Germany, 2014. Huynh, H. V., Ed.; In The Organometallic Chemistry of N-heterocyclic Carbenes; Wiley-VCH: Weinheim, Germany, 2017. Biju, A. T., Ed.; In N-Heterocyclic Carbenes in Organocatalysis; Wiley-VCH: Weinheim: Germany, 2018. Lappert, M. F. J. Organomet. Chem. 1988, 358, 185. Herrmann, W. A.; Köcher, C. Angew. Chem., Int. Ed. Engl. 1997, 36, 2162–2187. Arduengo, A. J., III. Acc. Chem. Res. 1999, 32, 913–921. Bourissou, D.; Guerret, O.; Gabbaï, F. P.; Bertrand, G. Chem. Rev. 2000, 100, 39–92. Herrmann, W. A. Angew. Chem. Int. Ed. 2002, 41, 1290–1309. Crabtree, R. H. J. Organomet. Chem. 2005, 690, 5451–5457. Hahn, F. E.; Jahnke, M. C. Angew. Chem. Int. Ed. 2008, 47, 3122–3172. Díez-González, S.; Marion, N.; Nolan, S. P. Chem. Rev. 2009, 109, 3612–3676. Hopkinson, M. N.; Richter, C.; Schedler, M.; Glorius, F. Nature 2014, 510, 485–496. Melaimi, M.; Jazzar, R.; Soleilhavoup, M.; Bertrand, G. Angew. Chem. Int. Ed. 2017, 56, 10046–10068. Soleilhavoup, M.; Bertrand, G. Chem 2020, 6, 1275–1282. Bézier, D.; Sortais, J.-B.; Darcel, C. Adv. Synth. Catal. 2013, 355, 19. Marion, N.; Nolan, S. P. Chem. Soc. Rev. 2008, 37, 1776–1782. Würtz, S.; Glorius, F. Acc. Chem. Res. 2008, 41, 1523–1533. Samojłowicz, C.; Bieniek, M.; Grela, K. Chem. Rev. 2009, 109, 3708–3742. Lin, J. C. Y.; Huang, R. T. W.; Lee, C. S.; Bhattacharyya, A.; Hwang, W. S.; Lin, I. J. B. Chem. Rev. 2009, 109, 3561–3598. Arnold, P. L.; Casely, I. J. Chem. Rev. 2009, 109, 3599–3611. Vougioukalakis, G. C.; Grubbs, R. H. Chem. Rev. 2010, 110, 1746–1787. Gil, W.; Trzeciak, A. M. Coord. Chem. Rev. 2011, 255, 473–483. Fortman, G. C.; Nolan, S. P. Chem. Soc. Rev. 2011, 40, 5151–5169.

Organometallic Chemistry of NHCs and Analogues 30. 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. 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.

361

Wang, F.; Liu, L.-J.; Wang, W.; Qu, M.; Zhao, M.-X.; Liu, L.-J.; Shi, M. Coord. Chem. Rev. 2012, 256, 804–853. Hock, S. J.; Schaper, L.-A.; Herrmann, W. A.; Kühn, F. E. Chem. Soc. Rev. 2013, 42, 5073–5089. Crabtree, R. H. Coord. Chem. Rev. 2013, 257, 755–766. Bellemin-Laponnaz, S.; Dagorne, S. Chem. Rev. 2014, 114, 8747–8774. Riener, K.; Haslinger, S.; Raba, A.; Högerl, M. P.; Cokoja, M.; Herrmann, W. A.; Kühn, F. E. Chem. Rev. 2014, 114, 5215–5272. Zhang, D.; Zi, G. Chem. Soc. Rev. 2015, 44, 1898–1921. Lazreg, F.; Nahra, F.; Cazin, C. S. J. Coord. Chem. Rev. 2015, 293–294, 48–79. Wang, Z.; Jiang, L.; Mohamed, D. K. B.; Zhao, J.; Hor, T. S. A. Coord. Chem. Rev. 2015, 293− 294, 292–326. Cheng, J.; Wang, L.; Wang, P.; Deng, L. Chem. Rev. 2018, 118, 9930–9987. Peris, E. Chem. Rev. 2018, 118, 9988–10031. Sipos, G.; Dorta, R. Coord. Chem. Rev. 2018, 375, 13–68. Danopoulos, A. A.; Simler, T.; Braunstein, P. Chem. Rev. 2019, 119, 3730–3961. Liang, Q.; Song, D. Chem. Soc. Rev. 2020, 49, 1209–1232. Zhao, Q.; Meng, G.; Nolan, S. P.; Szostak, M. Chem. Rev. 2020, 120, 1981–2048. Schuster, O.; Yang, L. R.; Raubenheimer, H. G.; Albrecht, M. Chem. Rev. 2009, 109, 3445–3478. Benhamou, L.; Chardon, E.; Lavigne, G.; Bellemin-Laponnaz, S.; César, V. Chem. Rev. 2011, 111, 2705–2733. Martin, D.; Melaimi, M.; Soleilhavoup, M.; Bertrand, G. Organometallics 2011, 30, 5304–5313. Soleilhavoup, M.; Bertrand, G. Acc. Chem. Res. 2015, 48, 256–266. Moerdyk, J. P.; Schilter, D.; Bielawski, C. W. Acc. Chem. Res. 2016, 49, 1458–1468. Vivancos, Á.; Segarra, C.; Albrecht, M. Chem. Rev. 2018, 118, 9493–9586. Guisado-Barrios, G.; Soleilhavoup, M.; Bertrand, G. Acc. Chem. Res. 2018, 51, 3236–3244. Sau, S. C.; Hota, P. K.; Mandal, S. K.; Soleilhavoup, M.; Bertrand, G. Chem. Soc. Rev. 2020, 49, 1233–1252. Huynh, H. V. Chem. Rev. 2018, 118, 9457–9492. Nelson, D. J.; Nolan, S. P. Chem. Soc. Rev. 2013, 42, 6723–6753. Bernhammer, J. C.; Frison, G.; Huynh, H. V. Chem. Eur. J. 2013, 19, 12892–12905. Jacobsen, H.; Correa, A.; Poater, A.; Costabile, C.; Cavallo, L. Coord. Chem. Rev. 2009, 253, 687–703. Comas-Vives, A.; Harvey, J. N. Eur. J. Inorg. Chem. 2011, 5025–5035. Jahnke, M. C.; Hahn, F. E. Coord. Chem. Rev. 2015, 293 − 294, 95–115. Perry, M. C.; Burgess, K. Tetrahedron: Asymmetry 2003, 14, 951–961. César, V.; Bellemin-Laponnaz, S.; Gade, L. H. Chem. Soc. Rev. 2004, 33, 619–636. Janssen-Müller, D.; Schlepphorst, C.; Glorius, F. Chem. Soc. Rev. 2017, 46, 4845–4854. Hahn, F. E. ChemCatChem 2013, 5, 419–430. Schaper, L.-A.; Hock, S. J.; Herrmann, W. A.; Kühn, F. E. Angew. Chem. Int. Ed. 2013, 52, 270–289. Jahnke, M. C.; Hahn, F. E. Chem. Lett. 2015, 44, 226–237. Kuwata, S.; Hahn, F. E. Chem. Rev. 2018, 118, 9642–9677. Peris, E.; Crabtree, R. H. Coord. Chem. Rev. 2004, 248, 2239–2246. Peris, E. Top. Organomet. Chem. 2007, 21, 83–116. Mata, J. A.; Poyatos, M.; Peris, E. Coord. Chem. Rev. 2007, 251, 841–859. Pugh, D.; Danopoulos, A. A. Coord. Chem. Rev. 2007, 251, 610–641. Poyatos, M.; Mata, J. A.; Peris, E. Chem. Rev. 2009, 109, 3677–3707. Biffis, A.; Baron, M.; Tubaro, C. Poly-NHC complexes of transition-metals: recent applications and new trends. In Adv. Organomet. Chem. Pérez, P. J., Ed.; Academic Press, 2015; pp 203–288. De, S.; Udvardy, A.; Czégéni, E. C.; Joó, F. Coord. Chem. Rev. 2019, 400, 1–31. Kühl, O. Chem. Soc. Rev. 2007, 36, 592–607. Liddle, S. T.; Edworthy, I. S.; Arnold, P. L. Chem. Soc. Rev. 2007, 36, 1732. Normand, A. T.; Cavell, K. J. Eur. J. Inorg. Chem. 2008, 2781–2800. Mo, Z.; Deng, L. Synlett 2014, 25, 1045–1049. Hameury, S.; de Fremont, P.; Braunstein, P. Chem. Soc. Rev. 2017, 46, 632–733. Fliedel, C.; Braunstein, P. J. Organomet. Chem. 2014, 751, 286–300. Tamm, M.; Hahn, F. E. Coord. Chem. Rev. 1999, 182, 175–209. Garrison, J. C.; Youngs, W. J. Chem. Rev. 2005, 105, 3978–4008. Hindi, K. M.; Panzner, M. J.; Tessier, C. A.; Cannon, C. L.; Youngs, W. J. Chem. Rev. 2009, 109, 3859–3884. Cronje, S.; Raubenheimer, H. G. Chem. Soc. Rev. 2008, 37, 1998–2011. Barnard, P. J.; Berners-Price, S. J. Coord. Chem. Rev. 2007, 251, 1889–1902. Visbal, R.; Concepción Gimeno, M. Chem. Soc. Rev. 2014, 43, 3551–3574. Sinha, N.; Hahn, F. E. Acc. Chem. Res. 2017, 50, 2167–2184. Gan, M.-M.; Liu, J.-Q.; Zhang, L.; Wang, Y.-Y.; Hahn, F. E.; Han, Y.-F. Chem. Rev. 2018, 118, 9587–9641. Nesterov, V.; Reiter, D.; Bag, P.; Frisch, P.; Holzner, R.; Porzelt, A.; Inoue, S. Chem. Rev. 2018, 118, 9678–9842. Moss, G. P.; Smith, P. A. S.; Tavernier, D. Pure Appl. Chem. 1995, 67, 1307–1375. Gleiter, R.; Hoffmann, R. J. Am. Chem. Soc. 1968, 90, 1475–1485. Kirmse, W., Ed.; In Carbene Chemistry, 2nd ed. Academic Press: New York, 1971. Moss, R. A., Jones, M., Jr., Eds.; In Carbenes, Vol. I; Wiley: New York, 1975. 1973, Vol. II. Advances in Carbene Chemistry, (Ed. Brinker, U.), Vol. I, JAI Press, Greenwich, (CT), 1994, Vol. II, JAI Press, Stamford (CT), 1998, Vol. III, Elsevier, Amsterdam, 2001. Bertrand, G., Ed.; In Carbene Chemistry; Marcel Dekker: New York, 2002. Harrison, J. F. J. Am. Chem. Soc. 1971, 93, 4112–4119. Bauschlicher, C. W., Jr.; Schaefer, H. F., III; Bagus, P. S. J. Am. Chem. Soc. 1977, 99, 7106–7110. Harrison, J. F.; Liedtke, R. C.; Liebman, J. F. J. Am. Chem. Soc. 1979, 101, 7162–7168. Feller, D.; Borden, W. T.; Davidson, E. R. Chem. Phys. Lett. 1980, 71, 22–26. Moss, R. A., Platz, M. S., Jones, M., Jr., Eds.; In Reactive Intermediate Chemistry; Wiley-Interscience: Hoboken, NJ, 2004. Dumas, J. B.; Péligot, E. Ann. Chim. Phys. 1835, 58, 5–74. Geuther, A. Ann. Chem. Pharm. 1862, 123, 121–122. Nef, J. U. Justus Liebigs Ann. Chem. 1895, 287, 265–359. Nef, J. U. Justus Liebigs Ann. Chem. 1897, 298, 202–374.

362 102. 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. 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.

Organometallic Chemistry of NHCs and Analogues Doering, W. V. E.; Hoffmann, A. K. J. Am. Chem. Soc. 1954, 76, 6162–6165. Breslow, R. J. Am. Chem. Soc. 1958, 80, 3719–3726. Moss, R. A. Acc. Chem. Res. 1989, 22, 15. Igau, A.; Grützmacher, H.; Baceiredo, A.; Bertrand, G. J. Am. Chem. Soc. 1988, 110, 6463–6466. Gillette, G.; Baceiredo, A.; Bertrand, G. Angew. Chem. Int. Ed. Engl. 1990, 29, 1429–1431. Kato, T.; Gornitzka, H.; Baceiredo, A.; Savin, A.; Bertrand, G. J. Am. Chem. Soc. 2000, 122, 998–999. Wanzlick, H.-W.; Schikora, E. Angew. Chem. 1960, 72, 494. Öfele, K. J. Organomet. Chem. 1968, 12, 42–43. Wanzlick, H.-W. Angew. Chem. Int. Ed. Engl. 1962, 1, 75–80. Wanzlick, H. W.; Schonher, H. J. Angew. Chem. Int. Ed. Engl. 1968, 7, 141–142. Cardin, D. J.; Cetinkaya, B.; Lappert, M. F.; Manojlovic-Muir, L.; Muir, K. W. Chem. Commun. 1971, 400–401. Hudnall, T. W.; Moorhead, E. J.; Gusev, D. G.; Bielawski, C. W. J. Org. Chem. 2010, 75, 2763–2766. Hudnall, T. W.; Bielawski, C. W. J. Am. Chem. Soc. 2009, 131, 16039–16041. César, V.; Lugan, N.; Lavigne, G. Eur. J. Inorg. Chem. 2010, 2010, 361–365. Hudnall, T. W.; Moerdyk, J. P.; Bielawski, C. W. Chem. Commun. 2010, 46, 4288–4290. Arduengo, A. J., III; Goerlich, J. R.; Marshall, W. J. Liebig’s Ann. 1997, 365–374. Lavallo, V.; Canac, Y.; Pr-sang, C.; Donnadieu, B.; Bertrand, G. Angew. Chem. Int. Ed. 2005, 44, 5705–5709. Jazzar, R.; Dewhurst, R. D.; Bourg, J. B.; Donnadieu, B.; Canac, Y.; Bertrand, G. Angew. Chem. Int. Ed. 2007, 46, 2899–2902. Jazzar, R.; Bourg, J. B.; Dewhurst, R. D.; Donnadieu, B.; Bertrand, G. R. J. Org. Chem. 2007, 72, 3492–3499. Hu, X.; Soleilhavoup, M.; Melaimi, M.; Chu, J.; Bertrand, G. Angew. Chem. Int. Ed. 2015, 54, 6008–6011. Aldeco-Perez, E.; Rosenthal, A. J.; Donnadieu, B.; Parameswaran, P.; Frenking, G.; Bertrand, G. Science 2009, 326, 556–559. Gründemann, S.; Kovacevic, A.; Albrecht, M.; Faller, J. W.; Crabtree, R. H. Chem. Commun. 2001, 2274–2275. Lebel, H.; Janes, M. K.; Charette, A. B.; Nolan, S. P. J. Am. Chem. Soc. 2004, 126, 5046–5047. Heckenroth, M.; Kluser, E.; Neels, A.; Albrecht, M. Angew. Chem. Int. Ed. 2007, 46, 6293–6296. Mathew, P.; Neels, A.; Albrecht, M. J. Am. Chem. Soc. 2008, 130, 13534–13535. Guisado-Barrios, G.; Bouffard, J.; Donnadieu, B.; Bertrand, G. Angew. Chem. Int. Ed. 2010, 49, 4759–4762. Ghadwal, R. S. Dalon. Trans. 2016, 45, 16081–16095. Jafarpour, L.; Stevens, E. D.; Nolan, S. P. J. Organomet. Chem. 2000, 606, 49–54. Bantreil, X.; Nolan, S. P. Nat. Protocols. 2011, 6, 69–77. Kuhn, N.; Kratz, T. Synthesis 1993, 561–562. Enders, D.; Breuer, K.; Raabe, G.; Runsink, J.; Teles, J. H.; Melder, J.-P.; Ebel, K.; Brode, S. Angew. Chem. Int. Ed. Engl. 1995, 34, 1021–1023. Nyce, G. W.; Csihony, S.; Waymouth, R. M.; Hedrick, J. L. Chem. Eur. J. 2004, 10, 4073–4079. Bantu, B.; Pawar, G. M.; Decker, U.; Wurst, K.; Schmidt, A. M.; Buchmeiser, M. R. Chem. Eur. J. 2009, 15, 3103–3109. Otto, M.; Conejero, S.; Canac, Y.; Romanenko, V. D.; Rudzevitch, V.; Bertrand, G. J. Am. Chem. Soc. 2004, 126, 1016–1017. Arduengo, A. J.; Bock, H.; Chen, H.; Denk, M.; Dixon, D. A.; Green, J. C.; Herrmann, W. A.; Jones, N. L.; Wagner, M.; West, R. J. Am. Chem. Soc. 1994, 116, 6641–6649. Dixon, D. A.; Arduengo, A. J. J. Phys. Chem. 1991, 95, 4180–4182. Arduengo, A. J.; Dixon, D. A.; Kumashiro, K. K.; Lee, C.; Power, W. P.; Zilm, K. W. J. Am. Chem. Soc. 1994, 116, 6361–6367. Heinemann, C.; Thiel, W. Chem. Phys. Lett. 1994, 217, 11–16. Carter, E. A.; Goddard, W. A. J. Phys. Chem. 1986, 90, 998–1001. Itoh, T.; Nakata, Y.; Hirai, K.; Tomioka, H. J. Am. Chem. Soc. 2005, 128, 957–967. Heinemann, C.; Müller, T.; Apeloig, Y.; Schwarz, H. J. Am. Chem. Soc. 1996, 118, 2023–2038. Boehme, C.; Frenking, G. J. Am. Chem. Soc. 1996, 118, 2039–2046. Tolman, C. A. J. Am. Chem. Soc. 1970, 92, 2953–2956. Tolman, C. A. Chem. Rev. 1977, 77, 313–348. Lappert, M. F.; Pye, P. L. J. Chem. Soc., Dalton Trans. 1977, 2172–2180. Öfele, K.; Herrmann, W. A.; Mihalios, D.; Elison, M.; Herdtweck, E.; Scherer, W.; Mink, J. J. Organomet. Chem. 1993, 459, 177–184. Denk, K.; Sirsch, P.; Herrmann, W. A. J. Organomet. Chem. 2002, 649, 219–224. Chianese, A.; Li, X.; Janzen, M.; Faller, J.; Crabtree, R. Organometallics 2003, 22, 1663–1667. Dorta, R.; Stevens, E. D.; Hoff, C. D.; Nolan, S. P. J. Am. Chem. Soc. 2003, 125, 10490–10491. Dorta, R.; Stevens, E. D.; Scott, N. M.; Costabile, C.; Cavallo, L.; Hoff, C. D.; Nolan, S. P. J. Am. Chem. Soc. 2005, 127, 2485–2495. Herrmann, W. A.; Goossen, L. J.; Köcher, C.; Artus, G. R. J. Angew. Chem. Int. Ed. Engl. 1996, 35, 2805–2807. Benhamou, L.; Vujkovic, N.; César, V.; Gornitzka, H.; Lugan, N.; Lavigne, G. Organometallics 2010, 29, 2616–2630. Meiries, S.; Speck, K.; Cordes, D. B.; Slawin, A. M. Z.; Nolan, S. P. Organometallics 2012, 32, 330–339. Collado, A.; Balogh, J.; Meiries, S.; Slawin, A. M. Z.; Falivene, L.; Cavallo, L.; Nolan, S. P. Organometallics 2013, 32, 3249–3252. Meiries, S.; Le Duc, G.; Chartoire, A.; Collado, A.; Speck, K.; Arachchige, K. S. A.; Slawin, A. M. Z.; Nolan, S. P. Chem. Eur. J. 2013, 19, 17358–17368. Paul, U. S. D.; Sieck, C.; Haehnel, M.; Hammond, K.; Marder, T. B.; Radius, U. Chem. Eur. J. 2016, 22, 11005–11014. Paul, U. S. D.; Radius, U. Organometallics 2017, 36, 1398–1407. Savka, R.; Plenio, H. Dalton Trans. 2015, 44, 891–893. Herrmann, W. A.; Schütz, J.; Frey, G. D.; Herdtweck, E. Organometallics 2006, 25, 2437–2448. Leuthaüßer, S.; Schwarz, D.; Plenio, H. Chem. Eur. J. 2007, 13, 7195–7203. Fürstner, A.; Alcarazo, M.; Krause, H.; Lehmann, C. W. J. Am. Chem. Soc. 2007, 129, 12676–12677. Bittermann, A.; Härter, P.; Herdtweck, E.; Hoffmann, S. D.; Herrmann, W. A. J. Organomet. Chem. 2008, 693, 2079–2090. Wolf, S.; Plenio, H. J. Organomet. Chem. 2009, 694, 1487–1492. Dröge, T.; Glorius, F. Angew. Chem. Int. Ed. 2010, 49, 6940–6952. Schaper, L.-A.; Ofele, K.; Kadyrov, R.; Bechlars, B.; Drees, M.; Cokoja, M.; Herrmann, W. A.; Kuhn, F. E. Chem. Commun. 2012, 48, 3857–3859. Sanderson, M. D.; Kamplain, J. W.; Bielawski, C. W. J. Am. Chem. Soc. 2006, 128, 16514–16515. Buhl, H.; Ganter, C. J. Organomet. Chem. 2016, 809, 74–78. Lavallo, V.; Canac, Y.; DeHope, A.; Donnadieu, B.; Bertrand, G. Angew. Chem. Int. Ed. 2005, 44, 7236–7239. Iglesias, M.; Beetstra, D. J.; Kariuki, B.; Cavell, K. J.; Dervisi, A.; Fallis, I. A. Eur. J. Inorg. Chem. 2009, 1913–1919. Neveling, A.; Julius, G. R.; Cronje, S.; Esterhuysen, C.; Raubenheimer, H. G. Dalton Trans. 2005, 181–192. Ung, G.; Bertrand, G. Chem. Eur. J. 2011, 17, 8269–8272. Terashima, T.; Inomata, S.; Ogata, K.; Fukuzawa, S. Eur. J. Inorg. Chem. 2012, 1387–1393. Lever, A. B. P. Inorg. Chem. 1990, 29, 1271–1285.

Organometallic Chemistry of NHCs and Analogues 175. 176. 177. 178. 179. 180. 181. 182. 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. 233. 234. 235. 236. 237. 238. 239. 240. 241. 242. 243. 244. 245. 246. 247.

Lever, A. B. P. Inorg. Chem 1991, 30, 1980–1985. Fielder, S. S.; Osborne, M. C.; Lever, A. B. P.; Pietro, W. J. J. Am. Chem. Soc. 1995, 117, 6990–6993. Leuthäußer, S.; Schmidts, V.; Thiele, C. M.; Plenio, H. Chem. Eur. J. 2008, 14, 5465–5481. Savka, R. D.; Plenio, H. J. Organomet. Chem. 2012, 710, 68–74. Vorfalt, T.; Leuthäußer, S.; Plenio, H. Angew. Chem. Int. Ed. 2009, 48, 5191–5194. Moerdyk, J. P.; Bielawski, C. W. Organometallics 2011, 30, 2278–2284. Blake, G. A.; Moerdyk, J. P.; Bielawski, C. W. Organometallics 2012, 31, 3373–3378. Tapu, D.; Dixon, D. A.; Roe, C. Chem. Rev. 2009, 109, 3385–3407. Cardin, D. J.; Cetinkaya, B.; Cetinkaya, E.; Lappert, M. F.; Randall, E. W.; Rosenberg, E. J. Chem. Soc., Dalton Trans. 1973, 1982. Herrmann, W. A.; Runte, O.; Artus, G. J. Organomet. Chem. 1995, 501, C1–C4. Huynh, H. V.; Han, Y.; Jothibasu, R.; Yang, J. A. Organometallics 2009, 28, 5395–5404. Teng, Q.; Huynh, H. V. Dalton Trans. 2017, 46, 614–627. Guo, S.; Sivaram, H.; Yuan, D.; Huynh, H. V. Organometallics 2013, 32, 3685–3696. Jothibasu, R.; Huynh, H. V.; Koh, L. L. J. Organomet. Chem. 2008, 693, 374–380. Joost, M.; Estevez, L.; Mallet-Ladeira, S.; Miqueu, K.; Amgoune, A.; Bourissou, D. Angew. Chem. Int. Ed. 2014, 53, 14512–14516. Garrison, J. C.; Simons, R. S.; Kofron, W. G.; Tessier, C. A.; Youngs, W. J. Chem. Commun. 2001, 1780–1781. Hu, X.; Castro-Rodriguez, I.; Olsen, K.; Meyer, K. Organometallics 2004, 23, 755–764. Cavallo, L.; Correa, A.; Costabile, C.; Jacobsen, H. J. Organomet. Chem. 2005, 690, 5407–5413. Back, O.; Henry-Ellinger, M.; Martin, C. D.; Martin, D.; Bertrand, G. Angew. Chem. Int. Ed. 2013, 52, 2939–2943. Liske, A.; Verlinden, K.; Buhl, H.; Schaper, K.; Ganter, C. Organometallics 2013, 32, 5269–5272. Verlinden, K.; Buhl, H.; Frank, W.; Ganter, C. Eur. J. Inorg. Chem. 2015, 2015, 2416–2425. Rodrigues, R. R.; Dorsey, C. L.; Arceneaux, C. A.; Hudnall, T. W. Chem. Commun. 2014, 50, 162–164. Buhl, H.; Verlinden, K.; Ganter, C.; Novakovic, S. B.; Bogdanovic, G. A. Eur. J. Inorg. Chem. 2016, 3389–3395. Vummaleti, S. V. C.; Nelson, D. J.; Poater, A.; Gómez-SuÁrez, A.; Cordes, D. B.; Slawin, A. M. Z.; Nolan, S. P.; Cavallo, L. Chem. Sci. 2015, 6, 1895–1904. Nelson, D. J.; Nahra, F.; Patrick, S. R.; Cordes, D. B.; Slawin, A. M. Z.; Nolan, S. P. Organometallics 2014, 33, 3640–3645. Thie, C.; Bruhn, C.; Leibold, M.; Siemeling, U. Molecules 2017, 22, 1133–1154. Manjare, S. T.; Singh, H. B.; Butcher, R. J. Tetrahedron 2012, 68, 10561–10566. Gusev, D. G. Organometallics 2009, 28, 763–770. Shi, Q.; Thatcher, R. J.; Slattery, J.; Sauari, P. S.; Whitwood, A. C.; McGowan, P. C.; Douthwaite, R. E. Chem. Eur. J. 2009, 15, 11346–11360. Tonner, R.; Frenking, G. Organometallics 2009, 28, 3901–3905. Gusev, D. G. Organometallics 2009, 28, 6458–6461. Cramer, C. J.; Truhlar, D. G. Phys. Chem. Chem. Phys. 2009, 11, 10757–10816. Fey, N.; Ridgway, B. M.; Jover, J.; McMullin, C. L.; Harvey, J. N. Dalton Trans. 2011, 40, 11184–11191. Mathew, J.; Suresh, C. H. Inorg. Chem. 2010, 49, 4665–4669. Mathew, J.; Suresh, C. H. Organometallics 2011, 30, 3106–3112. Setiawan, D.; Kalescky, R.; Kraka, E.; Cremer, D. Inorg. Chem. 2016, 55, 2332–2344. Clavier, H.; Nolan, S. P. Chem. Commun. 2010, 46, 841–861. Gómez-Suárez, A.; Nelson, D. J.; Nolan, S. P. Chem. Commun. 2017, 53, 2650–2660. Hillier, A. C.; Sommer, W. J.; Yong, B. S.; Petersen, J. L.; Cavallo, L.; Nolan, S. P. Organometallics 2003, 22, 4322–4326. Poater, A.; Cosenza, B.; Correa, A.; Giudice, S.; Ragone, F.; Scarano, V.; Cavallo, L. Eur. J. Inorg. Chem. 2009, 1759–1766. Poater, A.; Ragone, F.; Mariz, R.; Dorta, R.; Cavallo, L. Chem. – Eur. J. 2010, 16, 14348–14353. Falivene, L.; Credendino, R.; Poater, A.; Petta, A.; Serra, L.; Oliva, R.; Scarano, V.; Cavallo, L. Organometallics 2016, 35, 2286–2293. Balogh, J.; Slawin, A. M. Z.; Nolan, S. P. Organometallics 2012, 31, 3259–3263. Nelson, D. J.; Collado, A.; Manzini, S.; Meiries, S.; Slawin, A. M. Z.; Cordes, D. B.; Nolan, S. P. Organometallics 2014, 33, 2048–2058. Patrick, S. R.; Collado, A.; Meiries, S.; Slawin, A. M. Z.; Nolan, S. P. J. Organomet. Chem. 2015, 775, 152–154. Fructos, M. R.; Belderrain, T. R.; de Frémont, P.; Scott, N. M.; Nolan, S. P.; Díaz-Requejo, M. M.; Pérez, P. J. Angew. Chem. Int. Ed. 2005, 44, 5284–5288. Gaillard, S.; Bantreil, X.; Slawin, A. M. Z.; Nolan, S. P. Dalton Trans. 2009, 6967–6971. Gaillard, S.; Slawin, A. M. Z.; Bonura, A. T.; Stevens, E. D.; Nolan, S. P. Organometallics 2010, 29, 394–402. de Frémont, P.; Scott, N. M.; Stevens, E. D.; Nolan, S. P. Organometallics 2005, 24, 2411–2418. Altenhoff, G.; Goddard, R.; Lehmann, C. W.; Glorius, F. J. Am. Chem. Soc. 2004, 126 (46), 15195–15201. Würtz, S.; Lohre, C.; Fröhlich, R.; Bergander, K.; Glorius, F. J. Am. Chem. Soc. 2009, 131, 8344–8345. Frey, G. D.; Dewhurst, R. D.; Kousar, S.; Donnadieu, B.; Bertrand, G. J. Organomet. Chem. 2008, 693, 1674–1682. Chu, J.; Munz, D.; Jazzar, R.; Melaimi, M.; Bertrand, G. J. Am. Chem. Soc. 2016, 138, 7884–7887. Ragone, F.; Poater, A.; Cavallo, L. J. Am. Chem. Soc. 2010, 132, 4249–4258. Liu, P.; Montgomery, J.; Houk, K. N. J. Am. Chem. Soc. 2011, 133, 6956–6959. Araki, K.; Kuwata, S.; Ikariya, T. Organometallics 2008, 27, 2176–2178. Jahnke, M. C.; Hahn, F. E. Top. Organomet. Chem. 2010, 30, 95–129. Kuwata, S.; Ikariya, T. Chem. Eur. J. 2011, 17, 3542–3556. Sundberg, R. J.; Shepherd, R. E.; Taube, H. J. Am. Chem. Soc. 1972, 94, 6558–6559. Sundberg, R. J.; Bryan, R. F.; Taylor, I. F.; Taube, H. J. Am. Chem. Soc. 1974, 96, 381–392. Ruiz, J.; Perandones, B. F. J. Am. Chem. Soc. 2007, 129, 9298–9299. Ruiz, J.; Berros, A.; Perandones, B. F.; Vivanco, M. Dalton Trans. 2009, 6999–7007. Ruiz, J.; Perandones, B. F. Chem. Commun. 2009, 2741–2743. Price, C.; Shipman, M.; Gummerson, S.; Houlton, A.; Clegg, W.; Elsegood, M. R. J. J. Chem. Soc., Dalton Trans. 2001, 353–354. Price, C.; Shipman, M. A.; Rees, N. H.; Elsegood, M. R. J.; Edwards, A. J.; Clegg, W.; Houlton, A. Chem. Eur. J. 2001, 7, 1194–1201. Galindo, M.; Houlton, A. Inorg. Chim. Acta 2009, 362, 625–633. Tan, K. L.; Bergman, R. G.; Ellman, J. A. J. Am. Chem. Soc. 2002, 124, 3202–3203. Miranda-Soto, V.; Grotjahn, D. B.; Cooksy, A. L.; Golen, J. A.; Moore, C. E.; Rheingold, A. L. Angew. Chem. Int. Ed. 2011, 50, 631–635. Toda, T.; Kuwata, S.; Ikariya, T. Chem. Eur. J. 2014, 20, 9539–9542. Toda, T.; Yoshinari, A.; Ikariya, T.; Kuwata, S. Chem. Eur. J. 2016, 22, 16675–16683. Brackemeyer, D.; Schulte to Brinke, C, ; Roelfes, F.; Hahn, F. E. Dalton Trans. 2017, 46, 4510–4513. Cepa, S.; Schulte to Brinke, C, ; Roelfes, F.; Hahn, F. E. Organometallics 2015, 34, 5454–5460. He, F.; Braunstein, P.; Wesolek, M.; Danopoulos, A. A. Chem. Commun. 2015, 51, 2814–2817.

363

364 248. 249. 250. 251. 252. 253. 254. 255. 256. 257. 258. 259. 260. 261. 262. 263. 264. 265. 266. 267. 268. 269. 270. 271. 272. 273. 274. 275. 276. 277. 278. 279. 280. 281. 282. 283. 284. 285. 286. 287. 288. 289. 290. 291. 292. 293. 294. 295. 296. 297. 298. 299. 300. 301. 302. 303. 304. 305. 306. 307. 308. 309. 310. 311. 312. 313. 314. 315. 316. 317. 318. 319.

Organometallic Chemistry of NHCs and Analogues He, F.; Wesolek, M.; Danopoulos, A. A.; Braunstein, P. Chem. Eur. J. 2016, 22, 2658–2671. Gomez-Lopez, J. L.; Chávez, D.; Parra-Hake, M.; Royappa, A. T.; Rheingold, A. L.; Grotjahn, D. B.; Miranda-Soto, V. Organometallics 2016, 35, 3148–3153. Marelius, D. C.; Darrow, E. H.; Moore, C. E.; Golen, J. A.; Rheingold, A. L.; Grotjahn, D. B. Chem. Eur. J. 2015, 21, 10988–10992. Hahn, F. E.; Langenhahn, V.; Lügger, T.; Pape, T.; Le Van, D. Angew. Chem. Int. Ed. 2005, 44, 3759–3763. Meier, N.; Hahn, F. E.; Pape, T.; Siering, C.; Waldvogel, S. R. Eur. J. Inorg. Chem. 2007, 1210–1214. Kaufhold, O.; Stasch, A.; Edwards, P. G.; Hahn, F. E. Chem. Commun. 2007, 1822–1824. Kaufhold, O.; Stasch, A.; Pape, T.; Hepp, A.; Edwards, P. G.; Newman, P. D.; Hahn, F. E. J. Am. Chem. Soc. 2009, 131, 306–317. Miranda-Soto, V.; Grotjahn, D. B.; DiPasquale, A. G.; Rheingold, A. L. J. Am. Chem. Soc. 2008, 130, 13200–13201. Hahn, F. E.; Naziruddin, A. R.; Hepp, A.; Pape, T. Organometallics 2010, 29, 5283–5288. Ikariya, T.; Bull, Chem. Soc. Jpn. 2011, 84, 1–16. Kuwata, S.; Ikariya, T. Chem. Commun. 2014, 50, 14290–14300. Norris, M. R.; Flowers, S. E.; Mathews, A. M.; Cossairt, B. M. Organometallics 2016, 35, 2778–2781. Mo, Z.; Xiao, J.; Gao, Y.; Deng, L. J. Am. Chem. Soc. 2014, 136, 17414–17417. Mo, Z.; Li, Y.; Lee, H.; Deng, L. Organometallics 2011, 30, 4687–4694. Gao, Y.; Li, G.; Deng, L. J. Am. Chem. Soc. 2018, 140, 2239–2250. Gao, Y.; Wang, L.; Deng, L. ACS Catal. 2018, 8, 9637–9646. Berthon-Gelloz, G.; Siegler, M. A.; Spek, A. L.; Tinant, B.; Reek, J. N. H.; Marko, I. E. Dalton Trans. 2010, 39, 1444–1446. Gómez-Suárez, A.; Ramón, R. S.; Songis, O.; Slawin, A. M. Z.; Cazin, C. S. J.; Nolan, S. P. Organometallics 2011, 30, 5463–5470. Meiries, S.; Chartoire, A.; Slawin, A. M. Z.; Nolan, S. P. Organometallics 2012, 31, 3402–3409. Dible, B. R.; Cowley, R. E.; Holland, P. L. Organometallics 2011, 30, 5123–5132. Sübner, M.; Plenio, H. Chem. Commun. 2005, 5417–5419. Endo, K.; Grubbs, R. H. J. Am. Chem. Soc. 2011, 133, 8525–8527. Ablialimov, O.; Ke¸dziorek, M.; Torborg, C.; Malinska, M.; Woz´niak, K.; Grela, K. Organometallics 2012, 31, 7316–7319. Lian, X.; Chen, W.; Dang, L.; Li, Y.; Ho, C.-Y. Angew. Chem. Int. Ed. 2017, 56, 9048–9052. Herrmann, W. A.; Goossen, L. J.; Köcher, C.; Artus, G. R. J. Angew. Chem. Int. Ed. 1996, 35, 2805–2807. Enders, D.; Gielen, H.; Raabe, G.; Runsink, J.; Teles, J. H. Chem. Ber. 1996, 129, 1483–1488. Enders, D.; Gielen, H.; Beuer, K. Tetrahedron: Asymmetry 1997, 8, 3571–3574. Kündig, E. P.; Seidel, T. M.; Jia, Y. X.; Bernardinelli, G. Angew. Chem. Int. Ed. 2007, 46, 8484–8487. Bolm, C.; Kesselgruber, M.; Raabe, G. Organometallics 2002, 21, 707–710. Ma, Y.; Song, C.; Ma, C.; Sun, Z.; Chai, Q.; Andrus, M. B. Angew. Chem. Int. Ed. 2003, 42, 5871–5874. Albright, A.; Gawley, R. E. J. Am. Chem. Soc. 2011, 133, 19680–19683. Cai, Y.; Yang, X.-T.; Zhang, S. Q.; Li, F.; Li, Y.-Q.; Ruan, L.-X.; Hong, X.; Shi, S.-L. Angew. Chem. Int. Ed. 2018, 57, 1376–1380. Van Veldhuizen, J. J.; Garber, S. B.; Kingsbury, J. S.; Hoveyda, A. H. J. Am. Chem. Soc. 2002, 124, 4954–4955. Knight, R. L.; Leeper, F. J. J. Chem. Soc. Perkin Trans. 1998, 1, 1891–1894. Kerr, M. S.; de Alaniz, J. R.; Rovis, T. J. Am. Chem. Soc. 2002, 124, 10298–10299. Chen, Y.; Pan, Y.; He, Y.-M.; Fan, Q.-H. Angew. Chem. Int. Ed. 2019, 58, 16831–16834. Yoshida, K.; Kamimura, T.; Kuwabara, H.; Yanagisawa, A. Chem. Commun. 2015, 51, 15442–15445. Glorius, F.; Altenhoff, G.; Goddard, R.; Lehmann, C. Chem. Commun 2002, 2704–2705. Despagnet-Ayoub, E.; Grubbs, R. H. J. Am. Chem. Soc. 2004, 126, 10198–10199. Alder, R. W.; Blake, M. E.; Bortolotti, C.; Bufali, S.; Butts, C. P.; Linehan, E.; Oliva, J. M.; Orpen, A. G.; Quayle, M. J. Chem. Commun. 1999, 241–242. Bazinet, P.; Yap, G. P. A.; Richeson, D. S. J. Am. Chem. Soc. 2003, 125, 13314–13315. Iglesias, M.; Beetstra, D. J.; Stasch, A.; Horton, P. N.; Hursthouse, M. B.; Coles, S. J.; Cavell, K. J.; Dervisi, A.; Fallis, I. A. Organometallics 2007, 26, 4800–4809. Jazzar, R.; Liang, H.; Donnadieu, B.; Bertrand, G. J. Organomet. Chem. 2006, 691, 3201–3205. Iglesias, M.; Beetstra, D. J.; Knight, J. C.; Ooi, L.-L.; Stasch, A.; Coles, S.; Male, L.; Hursthouse, M. B.; Cavell, K. J.; Dervisi, A.; Fallis, I. A. Organometallics 2008, 27, 3279–3289. Dunsford, J. J.; Tromp, D. S.; Cavell, K. J.; Elsevier, C. J.; Kariuki, B. M. Dalton Trans. 2013, 42, 7318–7329. Lu, W. Y.; Cavell, K. J.; Wixey, J. S.; Kariuki, B. Organometallics 2011, 30, 5649–5655. Chesnokov, G. A.; Topchiy, M. A.; Dzhevakov, P. B.; Gribanov, P. S.; Tukov, A. A.; Khrustalev, V. N.; Asachenko, A. F.; Nechaev, M. S. Dalton Trans. 2017, 46, 4331–4345. Cervantes-Reyes, A.; Rominger, F.; Rudolph, M.; Hashmi, A. S. K. Chem. Eur. J. 2019, 25, 11745–11757. Arduengo, A. J., III; Davidson, F.; Rasika Dias, H. V.; Goerlich, J. R.; Khasnis, D.; Marshall, W. J.; Prakasha, T. J. Am. Chem. Soc. 1997, 119, 12742–12749. Lee, P. H.-M.; Ko, C.-C.; Zhu, N.; Yam, V. W.-W. J. Am. Chem. Soc. 2007, 129, 6058–6059. Neilson, B. M.; Lynch, V. M.; Bielawski, C. W. Angew. Chem. Int. Ed. 2011, 50, 10322–10326. Neilson, B. M.; Bielawski, C. W. J. Am. Chem. Soc. 2012, 134, 12693–12699. Teator, A. J.; Tian, Y.; Chen, M.; Lee, J. K.; Bielawski, C. W. Angew.Chem. Int.Ed. 2015, 54, 11559–11563. Moerdyk, J. P.; Bielawski, C. W. Chem. Commun. 2014, 50, 4551–4553. Nasr, A.; Winkler, A.; Tamm, M. Coord. Chem. Rev. 2016, 316, 68–124. Kronig, S.; Theuergarten, E.; Daniliuc, C. G.; Jones, P. G.; Tamm, M. Angew. Chem. Int. Ed. 2012, 51, 3240–3244. Luan, X.; Mariz, R.; Robert, C.; Gatti, M.; Blumentritt, S.; Linden, A.; Dorta, R. Org. Lett. 2008, 10, 5569–5572. Kumar, R.; Tamai, E.; Ohnishi, A.; Nishimura, A.; Hoshimoto, Y.; Ohashi, M.; Ogoshi, S. Synthesis 2016, 48, 2789–2794. Gao, P.; Sipos, G.; Foster, D.; Dorta, R. ACS Catal. 2017, 7, 6060–6064. Ho, C.-Y.; Chan, C.-W.; He, L. Angew. Chem. Int. Ed. 2015, 54, 4512–4516. Jana, A.; Trzybinski, D.; Woz´niak, K.; Grela, K. Chem. Eur. J. 2018, 24, 891–897. Liu, L.; Ishida, N.; Ashida, S.; Murakami, M. Org. Lett. 2011, 13, 1666–1669. Park, J. K.; Lackey, H. H.; Rexford, M. D.; Kovnir, K.; Shatruk, M.; McQuade, D. T. Org. Lett. 2010, 12, 5008–5011. Martin, D.; Lassauque, N.; Donnadieu, B.; Bertrand, G. Angew. Chem. Int. Ed. 2012, 51, 6172–6175. Khramov, D. M.; Rosen, E. L.; Lynch, V. M.; Bielawski, C. W. Angew. Chem. Int. Ed. 2008, 47, 2267–2270. Siemeling, U.; Färber, C.; Bruhn, C. Chem. Commun. 2009, 98–100. Petrov, A. R.; Derheim, A.; Oetzel, J.; Leibold, M.; Bruhn, C.; Scheerer, S.; Oßwald, S.; Winter, R. F.; Siemeling, U, Inorg. Chem. 2015, 54, 6657–6670. Öfele, K. Angew. Chem. Int. Ed. Engl. 1969, 8, 916–917. Wanzlick, H.-W.; Kleiner, H.-J.; Lasch, I.; Füldner, H. U.; Steinmaus, H. Liebigs Ann. Chem. 1967, 708, 155–169. Cardin, D. J.; Çetinkaya, B.; Çetinkaya, E.; Lappert, M. F.; Manojlovic-Muir, L.; Muir, K. W. J. Organomet. Chem. 1972, 44, C59–C62. Yen, S. K.; Koh, L. L.; Hahn, F. E.; Huynh, H. V.; Hor, T. S. A. Organometallics 2006, 25, 5105–5112. Fehlhammer, W. P.; Bartel, K.; Völkl, A.; Achatz, D. Z. Naturforsch. 1982, 37b, 1044–1053.

Organometallic Chemistry of NHCs and Analogues 320. 321. 322. 323. 324. 325. 326. 327. 328. 329. 330. 331. 332. 333. 334. 335. 336. 337. 338. 339. 340. 341. 342. 343. 344. 345. 346. 347. 348. 349. 350. 351. 352. 353. 354. 355. 356. 357. 358. 359. 360. 361. 362. 363. 364. 365. 366. 367. 368. 369. 370. 371. 372. 373. 374. 375. 376. 377. 378. 379. 380. 381. 382. 383. 384. 385. 386. 387. 388. 389. 390.

365

Ruiz, J.; Perandones, B. F.; García, G.; Mosquera, M. E. G. Organometallics 2007, 26, 5687–5694. Tubaro, C.; Biffis, A.; Basato, M.; Benetollo, F.; Cavell, K. J.; Ooi, L.-L. Organometallics 2005, 24, 4153–4164. Lindner, R.; Wagner, C.; Steinborn, D. J. Am. Chem. Soc. 2009, 131, 8861–8874. Weinstein, C. M.; Junor, G. P.; Tolentino, D. R.; Jazzar, R. J.; Melaimi, M.; Bertrand, G. J. Am. Chem. Soc. 2018, 140, 9255–9260. McCarty, Z. R.; Lastovickova, D. N.; Bielawski, C. W. Chem. Commun. 2016, 52, 5447–5450. Zhou, Y.; Liu, Q.; Lv, W.; Pang, Q.; Ben, R.; Qian, Y.; Zhao, J. Organometallics 2013, 32, 3753–3759. Rao, B.; Tang, H.; Zeng, X.; Liu, L.; Melaimi, M.; Bertrand, G. Angew. Chem. Int. Ed. 2015, 54, 14915–14919. Manzano, R.; Rominger, F.; Hashmi, A. S. K. Organometallics 2013, 32, 2199–2203. Jazzar, R.; Soleilhavoup, M.; Bertrand, G. Chem. Rev. 2020, 120, 4141–4168. Scattolin, T.; Nolan, S. P. Trends Chem. 2020, 2, 721–736. Herrmann, W. A.; Schwarz, J.; Gardiner, M. G. Organometallics 1999, 18, 4082–4089. Douthwaite, R. E.; Haüssinger, D.; Green, M. L. H.; Silcock, P. J.; Gomes, P. T.; Martins, A. M.; Danopoulos, A. A. Organometallics 1999, 18, 4584–4590. Hahn, F. E.; Foth, M. J. Organomet. Chem. 1999, 585, 241–245. Mata, J. A.; Chianese, A. R.; Miecznikowski, J. R.; Poyatas, M.; Peris, E.; Faller, J. W.; Crabtree, R. H. Organometallics 2004, 23, 1253–1263. Miecznikowski, J. C.; Crabtree, R. H. Organometallics 2004, 23, 629–631. Hahn, F. E.; von Fehren, T.; Lügger, T. Inorg. Chim. Acta. 2005, 358, 4137–4144. Ahrens, S.; Herdtweck, E.; Goutal, S.; Strassner, T. Eur. J. Inorg. Chem. 2006, 1268–1274. Wang, J.-W.; Li, Q.-S.; Xu, F.-B.; Song, H.-B.; Zhang, Z. Z. Eur. J. Org. Chem. 2006, 1310–1316. Huynh, H. V.; Jothibasu, R. J. Organomet. Chem. 2011, 696, 3369–3375. Kernbach, U.; Ramm, M.; Luger, P.; Fehlhammer, W. P. Angew. Chem. Int. Ed. Engl. 1996, 35, 310–312. Fränkel, R.; Kernbach, U.; Bakola-Christianopoulou, M.; Plaia, U.; Suter, M.; Ponikwar, W.; NKth, H.; Moinet, C.; Fehlhammer, W. P. J. Organomet. Chem. 2001, 617–618, 530–545. Forshaw, A. P.; Bontchev, R. P.; Smith, J. M. Inorg. Chem. 2007, 46, 3792–3794. Hu, X. L.; Castro-Rodriguez, I.; Meyer, K. J. Am. Chem. Soc. 2003, 125, 12237–12245. Hu, X. L.; Castro-Rodriguez, I.; Meyer, K. Organometallics 2003, 22, 3016–3018. Hu, X. L.; Tang, Y. J.; Gantzel, P.; Meyer, K. Organometallics 2003, 22, 612–614. Dias, H. V. R.; Jin, W. C. Tetrahedron Lett. 1994, 35, 1365–1366. Shi, Z.; Thummel, R. P. J. Org. Chem. 1995, 60, 5935–5945. Baker, M. V.; Skelton, B. W.; White, A. H.; Williams, C. C. Organometallics 2002, 21, 2674–2678. Wong, W. W. H.; Vickers, M. S.; Cowley, A. R.; Paul, R. L.; Beer, P. D. Org. Biomol. Chem. 2005, 3, 4201–4208. Bass, H. M.; Cramer, S. A.; McCullough, A. S.; Bernstein, K. J.; Murdock, C. R.; Jenkins, D. M. Organometallics 2013, 32, 2160–2167. Ghavami, Z. S.; Anneser, M. R.; Kaiser, F.; Altmann, P. J.; Hofmann, B. J.; Schlagintweit, J. F.; Grivani, G.; Kühn, F. E. Chem. Sci. 2018, 9, 8307–8314. DeJesus, J. F.; Jenkins, D. M. Chem. Eur. J. 2020, 26, 1429–1435. Mageed, A. H. J. Organomet. Chem. 2019, 902, 120965–120987. Blom, B.; Tan, G.; Enthaler, S.; Inoue, S.; Epping, J. D.; Driess, M. J. Am. Chem. Soc. 2013, 135, 18108–18120. Duan, W.; Shi, M.; Rong, G. Chem. Commun. 2003, 2916–2917. Xu, Q.; Gu, P.; Jiang, H.; Wei, Y.; Shi, M. Chem. Rec. 2016, 16, 2736–2749. Liu, Y.; Shi, M.; Deng, L. Organometallics 2014, 33, 5660–5669. Nieto, I.; Bontchev, R. P.; Smith, J. M. Eur. J. Inorg. Chem. 2008, 2476–2480. Shishkov, I. V.; Rominger, F.; Hofmann, P. Organometallics 2009, 28, 3532–3536. Hickey, A. K.; Lee, W.-T.; Chen, C.-H.; Pink, M.; Smith, J. M. Organometallics 2016, 35, 3069–3073. Lutz, S. A.; Hickey, A. K.; Gao, Y.; Chen, C.-H.; Smith, J. M. J. Am. Chem. Soc. 2020, 142 (36), 15527–15535. Ye, S.; Kupper, C.; Meyer, S.; Andris, E.; Navratil, R.; Krahe, O.; Mondal, B.; Atanasov, M.; Bill, E.; Roithova, J.; Meyer, F.; Neese, F. J. Am. Chem. Soc. 2016, 138, 14312–14325. Findlay, N. J.; Park, S. R.; Schoenebeck, F.; Cahard, E.; Zhou, S.; Berlouis, L. E. A.; Spicer, M. D.; Tuttle, T.; Murphy, J. A. J. Am. Chem. Soc. 2010, 132, 15462–15464. Meyer, S.; Klawitter, I.; Demeshko, S.; Bill, E.; Meyer, F. Angew. Chem. Int. Ed. 2013, 52, 901–905. Cramer, S. A.; Jenkins, D. M. J. Am. Chem. Soc. 2011, 133, 19342–19345. Anneser, M. R.; Haslinger, S.; Pçthig, A.; Cokoja, M.; Basset, J.-M.; Kühn, F. E. Inorg. Chem. 2015, 54, 3797–3804. Chandrachud, P. P.; Bass, H. M.; Jenkins, D. M. Organometallics 2016, 35, 1652–1657. Cordes, C.; Morganti, M.; Klawitter, I.; Schremmer, C.; Dechert, S.; Meyer, F. Angew. Chem. Int. Ed. 2019, 58, 10855–10858. Boydston, A. J.; Williams, K. A.; Bielawski, C. W. J. Am. Chem. Soc. 2005, 127, 12496–12497. Boydston, A. J.; Rice, J. D.; Sanderson, M. D.; Dykhno, O. L.; Bielawski, C. W. Organometallics 2006, 25, 6087–6098. Khramov, D. M.; Boydston, A. J.; Bielawski, C. W. Angew. Chem. Int. Ed. 2006, 45, 6186–6189. Boydston, A. J.; Bielawski, C. W. Dalton Trans. 2006, 4073–4077. Baker, M. V.; Skelton, B. W.; White, A. H.; Williams, C. C. J. Chem. Soc., Dalton Trans. 2001, 111–120. Frank, M.; Maas, G.; Schatz, J. Eur. J. Org. Chem. 2004, 607–613. Hahn, F. E.; Radloff, C.; Pape, T.; Hepp, A. Organometallics 2008, 27, 6408–6410. Radloff, C.; Weigand, J. J.; Hahn, F. E. Dalton Trans. 2009, 9392–9394. Hahn, F. E.; Radloff, C.; Pape, T.; Hepp, A. Chem. Eur. J. 2008, 14, 10900–10904. Rit, A.; Pape, T.; Hahn, F. E. J. Am. Chem. Soc. 2010, 132, 4572–4573. Han, Y.-F.; Jin, G.-X.; Hahn, F. E. J. Am. Chem. Soc. 2013, 135, 9263–9266. Segarra, C.; Guisado-Barrios, G.; Hahn, F. E.; Peris, E. Organometallics 2014, 33, 5077–5080. Han, Y.-F.; Jin, G.-X.; Daniliuc, C. G.; Hahn, F. E. Angew. Chem. Int. Ed. 2015, 54, 4958–4962. Sinha, N.; Tan, T. T. Y.; Peris, E.; Hahn, F. E. Angew. Chem. Int. Ed. 2017, 56, 7393–7397. Sinha, N.; Stegemann, L.; Tan, T. T. Y.; Doltsinis, N. L.; Strassert, C. A.; Hahn, F. E. Angew. Chem. Int. Ed. 2017, 56, 2785–2789. Viciano, M.; Sanaú, M.; Peris, E. Organometallics 2007, 26, 6050–6054. Gonell, S.; Poyatos, M.; Peris, E. Chem. Eur. J. 2014, 20, 5746–5751. Martínez-Agramunt, V.; Ruiz-Botella, S.; Peris, E. Chem. Eur. J. 2017, 23, 6675–6681. Martínez-Agramunt, V.; Gusev, D. G.; Peris, E. Chem. Eur. J. 2018, 24, 14802–14807. Ibáñez, S.; Poyatos, M.; Peris, E. Angew. Chem. Int. Ed. 2017, 56, 9786–9790. Xiao, X.-Q.; Jin, G.-X. Dalton Trans. 2009, 9298–9303. Xiao, X.-Q.; Jia, A.-Q.; Lin, Y.-J.; Jin, G.-X. Organometallics 2010, 29, 4842–4848. Gan, M.-M.; Zhang, W.; Huo, X.-K.; Han, Y.-F. Sci. Sin Chim 2017, 47, 705–712.

366 391. 392. 393. 394. 395. 396. 397. 398. 399. 400. 401. 402. 403. 404. 405. 406. 407. 408. 409. 410. 411. 412. 413. 414. 415. 416. 417. 418. 419. 420. 421. 422. 423. 424. 425. 426. 427. 428. 429. 430. 431. 432. 433. 434. 435. 436. 437. 438. 439. 440. 441. 442. 443. 444. 445. 446. 447. 448. 449. 450. 451. 452. 453. 454. 455. 456. 457. 458. 459. 460. 461. 462. 463.

Organometallic Chemistry of NHCs and Analogues Zhang, L.; Han, Y.-F. Dalton Trans. 2018, 47, 4267–4272. Li, Y.; An, Y.-Y.; Fan, J.-Z.; Liu, X.-X.; Li, X.; Hahn, F. E.; Wang, Y.-Y.; Han, Y.-F. Angew. Chem. Int. Ed. 2020, 59, 10073–10080. Li, Y.; An, Y.-Y.; Fan, J.-Z.; Liu, X.-X.; Li, X.; Hahn, F. E.; Wang, Y.-Y.; Han, Y.-F. Angew. Chem. Int. Ed. 2019, 58, 2–10. Zhang, L.; Das, R.; Li, C.-T.; Wang, Y.-Y.; Hahn, F. E.; Han, Y.-F. Angew. Chem. Int. Ed. 2019, 58, 13360–13364. Ma, L.-L.; An, Y.-Y.; Sun, L.-Y.; Wang, Y.-Y.; Hahn, F. E.; Han, Y.-F. Angew. Chem. Int. Ed. 2019, 131, 4026–4031. Zhang, Y.-W.; Bai, S.; Wang, Y.-Y.; Han, Y.-F. J. Am. Chem. Soc. 2020, 142, 13614–13621. Wang, H. M. J.; Lin, I. J. B. Organometallics 1998, 17, 972–975. McGuinness, D. S.; Cavell, K. J. Organometallics 2000, 19, 741–748. Larocque, T. G.; Badaj, A. C.; Dastgir, S.; Lavoie, G. G. Dalton Trans. 2011, 40, 12705–12712. Nakano, R.; Nozaki, K. J. Am. Chem. Soc. 2015, 137, 10934–10937. Tao, W.; Nakano, R.; Ito, S.; Nozaki, K. Angew. Chem. Int. Ed. 2016, 55, 2835–2839. Mo, Z.; Liu, Y.; Deng, L. Angew. Chem. Int. Ed. 2013, 52, 10845–10849. Sun, J.; Luo, J.; Luo, Y.; Deng, L. Angew. Chem. Int. Ed. 2017, 56, 2720–2724. Yang, C.; Lee, H. M.; Nolan, S. P. Org. Lett. 2001, 3, 1511–1514. Holmes, J.; Pask, C. M.; Fox, M. A.; Willans, C. E. Chem. Commun. 2016, 52, 6443–6446. Simler, T.; Braunstein, P.; Danopoulos, A. A. Chem. Commun. 2016, 52, 2717–2720. Lee, H. M.; Chiu, P. L.; Zeng, J. Y. Inorg. Chim. Acta. 2004, 357, 4313–4321. Downing, S. P.; Danopoulos, A. A. Organometallics 2006, 25, 1337–1340. Danopoulos, A. A.; Winston, S.; Gelbrich, T.; Hursthousea, M. B.; Tooze, R. P. Chem. Commun. 2002, 482–483. Edworthy, I. S.; Rodden, M.; Mungur, S. A.; Davis, K. M.; Blake, A. J.; Wilson, C.; Schröder, M.; Arnold, P. L. J. Organomet. Chem. 2005, 690, 5710–5719. Gonzalez-Sebastian, L.; Chaplin, A. B. Inorg. Chim. Acta. 2017, 460, 22–28. Edworthy, I. S.; Blake, A. J.; Wilson, C.; Arnold, P. L. Organometallics 2007, 26, 3684–3689. Moser, M.; Wucher, B.; Kunz, D.; Rominger, F. Organometallics 2007, 26, 1024–1030. Danopoulos, A. A.; Tsoureas, N.; Wright, J. A.; Light, M. E. Organometallics 2004, 23, 166–168. Zhou, Y.; Chen, W. Organometallics 2007, 26, 2742–2746. Jeon, S.-J.; Waymouth, R. M. Dalton Trans. 2008, 437–439. Ye, J.; Jin, S.; Chen, W.; Qiu, H. Inorg. Chem. Commun. 2008, 11, 404–408. Scheele, U. J.; Dechert, S.; Meyer, F. Inorg. Chim. Acta. 2006, 359, 4891–4900. Liu, X.; Chen, W. Organometallics 2012, 31, 6614–6622. Aihara, H.; Matsuo, T.; Kawaguchi, H. Chem. Commun. 2003, 2204–2205. Lee, H. M.; Zeng, J. Y.; Hu, C.-H.; Lee, M.-T. Inorg. Chem. 2004, 43, 6822–6829. Catalano, V. J.; Malwitz, M. A.; Etogo, A. O. Inorg. Chem. 2004, 43, 5714–5724. Spencer, L. P.; Winston, S.; Fryzuk, M. D. Organometallics 2004, 23, 3372–3374. Hahn, F. E.; Jahnke, M. C.; Pape, T. Organometallics 2006, 25, 5927–5936. Chen, J. C. C.; Lin, I. J. B. J. Chem. Soc. Dalton Trans. 2000, 839–840. Peris, E.; Loch, J. A.; Mata, J.; Crabtree, R. H. Chem. Commun. 2001, 201–202. GrRndemann, S.; Albrecht, M.; Loch, J. A.; Faller, J. W.; Crabtree, R. H. Organometallics 2001, 20, 5485–5488. Danopoulos, A. A.; Winston, S.; Motherwell, W. B. Chem. Commun. 2002, 1376–1377. Danopoulos, A. A.; Tulloch, A. A. D.; Winston, S.; Eastham, G.; Hursthouse, M. B. Dalton Trans. 2003, 1009–1015. Miecznikowski, J. R.; Gründemann, S.; Albrecht, M.; Mégret, C.; Clot, E.; Faller, J. W.; Eisenstein, O.; Crabtree, R. H. Dalton Trans. 2003, 831–838. Douthwaite, R. E.; Houghton, J.; Kariuki, B. M. Chem. Commun. 2004, 698–699. Andavan, G. T. S.; Bauer, E. B.; Letko, C. S.; Hollis, T. K.; Tham, F. S. J. Organomet. Chem. 2005, 690, 5938–5947. Hahn, F. E.; Jahnke, M. C.; Gomez-Benitez, V.; Morales-Morales, D.; Pape, T. Organometallics 2005, 24, 6458–6463. Wright, J. A.; Danopoulos, A. A.; Motherwell, W. B.; Carroll, R. J.; Ellwood, S.; Saßmannshausen, J. Eur. J. Inorg. Chem. 2006, 4857–4865. Hahn, F. E.; Jahnke, M. C.; Pape, T. Organometallics 2007, 26, 150–154. Son, S. U.; Park, K. H.; Lee, Y. S.; Kim, B. Y.; Choi, C. H.; Lah, M. S.; Jang, Y. H.; Jang, D. J.; Chung, Y. K. Inorg. Chem. 2004, 43, 6896. Tu, T.; Malineni, J.; Dötz, K. H. Adv. Synth. Catal. 2008, 350, 1791–1795. Raynal, M.; Cazin, C. S. J.; Vallée, C.; Olivier-Bourbigou, H.; Braunstein, P. Chem. Commun. 2008, 3983–3985. Raynal, M.; Pattacini, R.; Cazin, C. S. J.; Vallée, C.; Olivier-Bourbigou, H.; Braunstein, P. Organometallics 2009, 28, 4028–4047. Huynh, H. V.; Yuan, D. Han, Y. Dalton Trans. 2009, ;7262–7268. Wucher, B.; Moser, M.; Schumacher, S. A.; Rominger, F.; Kunz, D. Angew. Chem. Int. Ed. 2009, 48, 4417–4421. Li, X. W.; Chen, F.; Xu, W. F.; Li, Y. Z.; Chen, X. T.; Xue, Z. L. Inorg. Chem. Commun. 2011, 14, 1673–1676. Yuan, D.; Tang, H.; Xiao, L.; Huynh, H. V. Dalton Trans. 2011, 40, 8788–8795. Huynh, H. V.; Lee, C.-S. Dalton Trans. 2013, 42, 6803–6809. Andrew, R. E.; Chaplin, A. B. Inorg. Chem. 2015, 54, 312–322. Seyboldt, A.; Wucher, B.; Alles, M.; Rominger, F.; Maichle-Mçssmer, C.; Kunz, D. J. Organomet. Chem. 2015, 775, 202–208. Chianese, A. R.; Mo, A.; Lampland, N. L.; Swartz, R. L.; Bremer, P. T. Organometallics 2010, 29, 3019–3026. Yu, R. P.; Darmon, J. M.; Hoyt, J. M.; Margulieux, G. W.; Turner, Z. R.; Chirik, P. J. ACS Catal. 2012, 2, 1760–1764. Xu, M.; Li, X.; Sun, Z.; Tu, T. Chem. Commun. 2013, 49, 11539–11541. Yu, R. P.; Darmon, J. M.; Milsmann, C.; Margulieux, G. W.; Stieber, S. C. E.; DeBeer, S.; Chirik, P. J. J. Am. Chem. Soc. 2013, 135, 13168–13184. Ibrahim, A. D.; Tokmic, K.; Brennan, M. R.; Kim, D.; Matson, E. M.; Nilges, M. J.; Bertke, J. A.; Fout, A. R. Dalton Trans. 2016, 45, 9805–9811. Tokmic, K.; Markus, C. R.; Zhu, L.; Fout, A. R. J. Am. Chem. Soc. 2016, 138, 11907–11913. Guisado-Barrios, G.; Bouffard, J.; Donnadieu, B.; Bertrand, G. Organometallics 2011, 30, 6017–6021. Poyatos, M.; McNamara, W.; Incarvito, C.; Peris, E.; Crabtree, R. H. Chem. Commun. 2007, 41, 2267–2269. Bezuidenhout, D. I.; Kleinhans, G.; Guisado-Barrios, G.; Liles, D. C.; Ung, G.; Bertrand, G. Chem. Commun. 2014, 50, 2431–2433. Iwasaki, H.; Yamada, Y.; Ishikawa, R.; Koga, Y.; Matsubara, K. Eur. J. Org. Chem. 2016, 1651–1654. Lee, W.-T.; Dickie, D. A.; Metta-Magaña, A. J.; Smith, J. M. Inorg. Chem. 2013, 52, 12842–12846. Andrada, D. M.; Holzmann, N.; Hamadi, T.; Frenking, G. Beilstein J. Org. Chem. 2015, 11, 2727–2736. Kleinhans, G.; Hansmann, M. M.; Guisado-Barrios, G.; Liles, D. C.; Bertrand, G.; Bezuidenhout, D. I. J. Am. Chem. Soc. 2016, 138, 15873–15876. Yan, X.; Bouffard, J.; Guisado-Barrios, G.; Donnadieu, B.; Bertrand, G. Chem. Eur. J. 2012, 18, 14627–14631. Suntrup, L.; Stein, F.; Klein, J.; Wilting, A.; Parlane, F. G. L.; Brown, C. M.; Fiedler, J.; Berlinguette, C. P.; Siewert, I.; Sarkar, B. Inorg. Chem. 2020, 59, 4215–4227. Maity, R.; van der Meer, M.; Hohloch, S.; Sarkar, B. Organometallics 2015, 34, 3090–3096. Hettmanczyk, L.; Suntrup, L.; Klenk, S.; Hoyer, C.; Sarkar, B. Chem. Eur. J. 2017, 23, 576–585.

Organometallic Chemistry of NHCs and Analogues 464. 465. 466. 467. 468. 469. 470. 471. 472. 473. 474. 475. 476. 477. 478. 479. 480. 481. 482. 483. 484. 485. 486. 487. 488. 489. 490. 491. 492. 493. 494. 495. 496. 497. 498. 499. 500. 501. 502. 503. 504. 505. 506. 507. 508. 509. 510. 511. 512. 513. 514. 515. 516. 517. 518. 519. 520. 521. 522. 523. 524. 525. 526. 527. 528. 529. 530. 531.

367

Jacobsen, H.; Ziegler, T. J. Am. Chem. Soc. 1994, 116, 3667–3679. Jacobsen, H.; Ziegler, T. Inorg. Chem. 1995, 17, 301–317. Boehme, C.; Frenking, G. Organometallics 1998, 17, 5801–5809. Tulloch, A. A. D.; Danopoulos, A. A.; Kleinhenz, S.; Light, M. E.; Hursthouse, M. B.; Eastham, G. Organometallics 2001, 20, 2027–2031. Nemcsok, D.; Wichmann, K.; Frenking, G. Organometallics 2004, 23, 3640–3646. Jacobsen, H. J. Organomet. Chem. 2005, 690, 6068–6078. Scott, N. M.; Dorta, R.; Stevens, E. D.; Correa, A.; Cavallo, L.; Nolan, S. P. J. Am. Chem. Soc. 2005, 127, 3516–3526. Jacobsen, H.; Correa, A.; Costabile, C.; Cavallo, L. J. Organomet. Chem. 2006, 691, 4350–4358. Tonner, R.; Heydenrych, G.; Frenking, G. Chem.-Asian J. 2007, 2, 1555–1567. Fantasia, S.; Petersen, J. L.; Jacobsen, H.; Cavallo, L.; Nolan, S. P. Organometallics 2007, 26, 5880–5889. Niehues, M.; Kehr, G.; Erker, G.; Wibbeling, B.; Fröhlich, R.; Blacque, O.; Berke, H. J. Organomet. Chem. 2002, 663, 192–203. Penka, E. F.; Schlaepfer, C. W.; Atanasov, M.; Albrecht, M.; Daul, C. J. Organomet. Chem. 2007, 692, 5709–5716. Ray, L.; Shaikh, M. M.; Ghosh, P. Dalton Trans. 2007, 4546–4555. Kausamo, A.; Tuononen, H. M.; Krahulic, K. E.; Roesler, R. Inorg. Chem. 2008, 47, 1145–1151. Díez-González, S.; Nolan, S. P. Coord. Chem. Rev. 2007, 251, 874–883. Radius, U.; Bickelhaupt, M. F. Organometallics 2008, 27, 3410–3414. Jin, L.; Tolentino, D. R.; Melaimi, M.; Bertrand, G. Sci. Adv. 2015, 1. No. e1500304. Jin, L.; Weinberger, D. S.; Melaimi, M.; Moore, C. E.; Rheingold, A. L.; Bertrand, G. Angew. Chem. Int. Ed. 2014, 53, 9059–9063. Weinberger, D. S.; Melaimi, M.; Moore, C. E.; Rheingold, A. L.; Frenking, G.; Jerabek, P.; Bertrand, G. Angew. Chem. Int. Ed. 2013, 52, 8964–8967. Weinberger, D. S.; Amin, S. K. N.; Mondal, K. C.; Melaimi, M.; Bertrand, G.; Stückl, A. C.; Roesky, H. W.; Dittrich, B.; Demeshko, S.; Schwederski, B.; Kaim, W.; Jerabek, P.; Frenking, G. J. Am. Chem. Soc. 2014, 136, 6235–6238. Ung, G.; Soleilhavoup, M.; Bertrand, G.; Peters, J. C. Angew. Chem. Int. Ed. 2014, 53, 8427–8431. Zeng, X.; Frey, G. D.; Kousar, S.; Bertrand, G. Chem. Eur. J. 2009, 15, 3056–3060. Samuel, P. P.; Mondal, K. C.; Roesky, H. W.; Hermann, M.; Frenking, G.; Demeshko, S.; Meyer, F.; Stückl, A. C.; Christian, J. H.; Dalal, N. S.; Ungur, L.; Chibotaru, L. F.; Pröpper, K.; Meents, A.; Dittrich, B. Angew. Chem. Int. Ed. 2013, 52, 11817–11821. Singh, A. P.; Samuel, P. P.; Roesky, H. W.; Schwarzer, M. C.; Frenking, G.; Sidhu, N. S.; Dittrich, B. J. Am. Chem. Soc. 2013, 135, 7324–7329. Zhang, L.; Liu, Y.; Deng, L. J. Am. Chem. Soc. 2014, 136, 15525–15528. Zhang, H.; Ouyang, Z.; Liu, Y.; Zhang, Q.; Wang, L.; Deng, L. Angew. Chem. Int. Ed. 2014, 53, 8432–8436. Roy, S.; Mondal, K. C.; Meyer, J.; Niepötter, B.; Köhler, C.; Herbst-Irmer, R.; Stalke, D.; Dittrich, B.; Andrada, D. M.; Frenking, G.; Roesky, H. W. Chem. Eur. J. 2015, 21, 9312–9318. Samuel, P. P.; Mondal, K. C.; Amin, S. K. N.; Roesky, H. W.; Carl, C.; Neufeld, R.; Stalke, D.; Demeshko, S.; Meyer, F.; Ungur, L.; Chibotaru, L. F.; Christian, J.; Ramachandran, V.; van Tol, J.; Dalal, N. S. J. Am. Chem. Soc. 2014, 136, 11964–11971. Arrowsmith, M.; Braunschweig, H.; Celik, M. A.; Dellermann, T.; Dewhurst, R. D.; Ewing, W. C.; Hammond, K.; Kramer, T.; Krummenacher, I.; Mies, J.; Radacki, K.; Schuster, J. K. Nat. Chem. 2016, 8, 890–894. Morokuma, K. J. Chem. Phys. 1971, 55, 1236–1244. Ziegler, T.; Rauk, A. Theor. Chim. Acta 1977, 46, 1–10. Ziegler, T.; Rauk, A. Inorg. Chem. 1979, 18, 1558–1565. Wagner, J. P.; Schreiner, P. R. J. Chem.Theory Comput. 2016, 12, 231–237. Weskamp, T.; Kohl, F. J.; Hieringer, W.; Gleich, D.; Herrmann, W. A. Angew. Chem. Int. Ed. 1999, 38, 2416–2419. Scholl, M.; Trnka, T. M.; Morgan, J. P.; Grubbs, R. H. Tetrahedron Lett. 1999, 40, 2247–2250. Huang, J.; Stevens, E. D.; Nolan, S. P.; Petersen, J. L. J. Am. Chem. Soc. 1999, 121, 2674–2678. Furst, M. R. L.; Cazin, C. S. J. Chem. Commun. 2010, 46, 6924–6925. Hitchcock, P. B.; Lappert, M. F.; Pye, P. L. J. Chem. Soc., Dalton Trans. 1977, 2160–2172. Lappert, M. F. J. Organomet. Chem. 2005, 690, 5467–5473. Fraser, P. J.; Roper, W. R.; Stone, F. G. A. J. Organomet. Chem. 1973, 50, C54–C56. McGuinness, D. S.; Cavell, K. J.; Yates, B. F. Chem. Commun. 2001, 355–356. McGuinness, D. S.; Cavell, K. J.; Yates, B. F.; Skelton, B. W.; White, A. H. J. Am. Chem. Soc. 2001, 123, 8317–8328. Marion, N.; de Frémont, P.; Puijk, I. M.; Ecarnot, E. C.; Amoroso, D.; Bell, A.; Nolana, S. P. Adv. Synth. Catal. 2007, 349, 2380–2384. Viciano, M.; Poyatos, M.; Sanaú, M.; Peris, E.; Rossin, A.; Ujaque, G.; Lledós, A. Organometallics 2006, 25 (5), 1120–1134. Holbrey, J. D.; Reichert, W. M.; Tkatchenko, I.; Bouajila, E.; Walter, O.; Tommasi, I.; Rogers, R. D. Chem. Commun. 2003, 28–29. Voutchkova, A. M.; Feliz, M.; Clot, E.; Eisenstein, O.; Crabtree, R. H. J. Am. Chem. Soc. 2007, 129, 12834–12846. Tommasi, I.; Sorrentino, F. Tetrahedron Lett. 2006, 47, 6453–6456. Févre, M.; Pinaud, J.; Leteneur, A.; Gnanou, Y.; Vignolle, J.; Taton, D.; Miqueu, K.; Sotiropoulos, J.-M. J. Am. Chem. Soc. 2012, 134, 6776–6784. Kösterke, T.; Kösters, J.; Würthwein, E.-U.; Mück-Lichtenfeld, C.; Schulte to Brinke, C, ; Lahoz, F.; Hahn, F. E. Chem. Eur. J. 2012, 18, 14594–14598. Jothibasu, R.; Huynh, H. V. Organometallics 2009, 28, 2505–2513. Liu, C.-Y.; Chen, D.-Y.; Lee, G.-H.; Peng, S.-M.; Liu, S.-T. Organometallics 1996, 15, 1055–1061. Hahn, F. E.; Imhof, L. Organometallics 1997, 16, 763–769. Hahn, F. E.; Tamm, M. J. Chem. Soc. Chem. Commun. 1995, 569–570. Hahn, F. E.; Langenhahn, V.; Meier, N.; Lügger, T.; Fehlhammer, W. P. Chem. – Eur. J. 2003, 9, 704–712. Dumke, A. C.; Pape, T.; Kösters, J.; Feldmann, K.-O.; Schulte to Brinke, C, ; Hahn, F. E. Organometallics 2013, 32, 289–299. Vogel, C.; Heinemann, F. W.; Sutter, J.; Anthon, C.; Meyer, K. Angew. Chem. Int. Ed. 2008, 47, 2681–2684. Scepaniak, J. J.; Vogel, C. S.; Khusniyarov, M. M.; Heinemann, F. W.; Meyer, K.; Smith, J. M. Science 2011, 331, 1049–1052. Kropp, H.; King, A. E.; Khusniyarov, M. M.; Heinemann, F. W.; Lancaster, K. M.; DeBeer, S.; Bill, E.; Meyer, K. J. Am. Chem. Soc. 2012, 134, 15538–15544. Goetz, M. K.; Anderson, J. S. J. Am. Chem. Soc. 2019, 141, 4051–4062. Martinez, J. L.; Lutz, S. A.; Xie, J.-Z.; Telser, J.; Hoffman, B. M.; Carta, V.; Pink, M.; Losovyj, Y.; Smith, M. J. Science 2020, 370, 356–359. Wang, L.; Hu, L.-R.; Zhang, H.-Z.; Chen, H.; Deng, L. J. Am. Chem. Soc. 2015, 137, 14196–14207. Laskowski, C. A.; Miller, A. J.; Hillhouse, G. L.; Cundari, T. R. J. Am. Chem. Soc. 2011, 133, 771–773. Du, J.-Z.; Wang, L.-B.; Xie, M.-H.; Deng, L. Angew. Chem. Int. Ed. 2015, 54, 12640–12644. Yao, X.-N.; Du, J.-Z.; Zhang, Y.-Q.; Leng, X.-B.; Yang, M.-W.; Jiang, S.-D.; Wang, Z.-X.; Ouyang, Z.-W.; Deng, L. J. Am. Chem. Soc. 2017, 139, 373–380. Cheng, J.; Liu, J.; Leng, X.-B.; Thomas, L.; Alexander, S.; Eckhard, B.; Ye, S.-F.; Deng, L. Inorg. Chem. 2019, 58, 7634–7644. Arduengo, A. J.; Camper, S. F.; Calabrese, J. C.; Davidson, F. J. Am. Chem. Soc. 1994, 116, 4391–4394. Hruszkewycz, D. P.; Wu, J.-G.; Hazari, N.; Incavito, C. D. J. Am. Chem. Soc. 2011, 133, 3280–3283. Fortman, G. C.; Scott, N. M.; Linden, A.; Stevens, E. D.; Dorta, R.; Nolan, S. P. Chem. Comm. 2010, 46, 1050–1052.

368 532. 533. 534. 535. 536. 537. 538. 539. 540. 541. 542. 543. 544. 545. 546. 547. 548. 549. 550. 551. 552. 553. 554. 555. 556. 557. 558. 559. 560. 561. 562. 563. 564. 565. 566. 567. 568. 569. 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.

Organometallic Chemistry of NHCs and Analogues Meng, Y.-S.; Mo, Z.-B.; Wang, B.-W.; Zhang, Y.-Q.; Deng, L.; Gao, S. Chem. Sci. 2015, 6, 7156–7162. Mo, Z.-B.; Ouyang, Z.-W.; Wang, L.; Fillman, K.; Neidig, M. L.; Deng, L. Org. Chem. Front. 2014, 1, 1040–1044. Ouyang, Z.-W.; Du, J.-Z.; Wang, L.; Kneebone, J. L.; Neidig, M. L.; Deng, L. Inorg. Chem. 2015, 54, 8808–8816. Ouyang, Z.; Meng, Y.; Cheng, J.; Xiao, J.; Gao, S.; Deng, L. Organometallics 2016, 35, 1361–1367. Jerabek, P.; Roesky, H. W.; Bertrand, G.; Frenking, G. J. Am. Chem. Soc. 2014, 136, 17123–17135. Mondal, K. C.; Samuel, P. P.; Li, Y.; Roy, S.; Roesky, H. W.; Ackermann, L.; Sidhu, N. S.; Sheldrick, G. M.; Carl, C.; Demeshko, S.; De, S.; Parameswaran, P.; Ungur, L.; Chibotaru, L. F.; Andrada, D. M. Eur. J. Inorg. Chem. 2014, 2014, 818–823. Du, J.-Z.; Chen, W.-W.; Chen, Q.; Leng, X.-B.; Meng, Y.-S.; Gao, S.; Deng, L. Organometallics 2020, 39, 729–739. Roy, S.; Mondal, K. C.; Mayer, J.; Niepotter, B.; Kohler, C.; Herbst-Irmer, R.; Stalke, D.; Dittrich, B.; Andrada, D. M.; Frenking, G.; Roesky, H. W. Chem. – Eur. J. 2015, 21, 9321–9328. Wu, J. G.; Faller, J. W.; Hazari, N.; Schmeier, T. J. Organometallics 2012, 31, 806–809. Jackstell, R.; Harkal, S.; Jiao, H.-J.; Spannenberg, A.; Borgmann, C.; Rottger, D.; Nierlich, F.; Elliot, M.; Niven, S.; Cavell, K.; Navarro, O.; Viciu, M. S.; Nolan, S. P.; Beller, M. Chem.-Eur. J. 2004, 10, 3891–3900. Marko, I. E.; Sterin, S.; Buisine, O.; Berthon, G.; Michaud, G.; Tinant, B.; Declercq, J. P. Adv Syn. Cata. 2004, 346, 1429–1434. Sun, J.; Gao, Y.-F.; Deng, L. Inorg. Chem. 2017, 56, 10775–10784. Wang, P.; Cheng, J.; Wang, D.-Y.; Yang, C.-B.; Leng, X.-B.; Deng, L. Organometallics 2020, 39, 2871–2877. Cheng, J.; Chen, Q.; Leng, X.-B.; Ye, S.-F.; Deng, L. Inorg. Chem. 2019, 58, 13129–13141. Cheng, J.; Chen, Q.; Leng, X.-B.; Ouyang, Z.-W.; Wang, Z.-X.; Ye, S.-F.; Deng, L. Chem 2018, 4, 2844–2860. Elsby, M. R.; Johnson, S. A. J. Am. Chem. Soc. 2017, 139, 9401–9407. Liu, T. B.; Darensbourg, M. Y. J. Am. Chem. Soc. 2007, 129, 7008–7009. Hsieh, C.-H.; Darensbourg, M. Y. J. Am. Chem. Soc. 2010, 132, 14118–14125. Deng, L.; Holm, R. H. J. Am. Chem. Soc. 2008, 130, 9878–9886. Brown, A. C.; Suess, D. M. J. Am. Chem. Soc. 2020, 142, 14240–14248. Evans, K. J.; Campbell, C. L.; Haddow, M. F.; Luz, C.; Morton, P. A.; Mansell, S. M. Eur. J. Inorg. Chem. 2019, 4894–4901. Lake, B. R. M.; Willans, C. E. Organometallics 2014, 33, 2027–2038. Catalano, V. J.; Munro, L. B.; Strasser, C. E.; Samin, A. F. Inorg. Chem. 2011, 50, 8465–8476. Thoi, V. S.; Chang, C. J. Chem. Commun. 2011, 47, 6578–6580. Herrmann, W. A.; Elison, M.; Fisher, J.; Köcher, C.; Artus, G. R. J. Angew. Chem. Int. Ed. Engl. 1995, 34, 2371–2373. Correa, A.; Nolan, S. P.; Cavallo, L. Top. Curr. Chem. 2011, 302, 131–155. Marion, N.; Nolan, S. P. Acc. Chem. Res. 2008, 41, 1440–1449. Jackson, E. P.; Malik, H. A.; Sormunen, G. J.; Baxter, R. D.; Liu, P.; Wang, H.; Shareef, A.-R.; Montgomery, J. Acc. Chem. Res. 2015, 48, 1736–1745. Hoshimoto, Y.; Ohashi, M.; Ogoshi, S. Acc. Chem. Res. 2015, 48, 1746–1755. Buchmeiser, M. R.; Sen, S.; Lienert, C.; Widmann, L.; Schowner, R.; Herz, K.; Hauser, P.; Frey, W.; Wang, D. ChemCatChem 2016, 8, 2710–2723. Schowner, R.; Elser, I.; Benedikter, M.; Momin, M.; Frey, W.; Schneck, T.; Stçhr, L.; Buchmeiser, M. R. Angew. Chem. Int. Ed. 2020, 59, 951–958. Kelly, R. A.; Scott, N. M.; Diez-Gonzalez, S.; Stevens, E. D.; Nolan, S. P. Organometallics 2005, 24, 3442. Mo, Z.; Zhang, Q.; Deng, L. Organometallics 2012, 31, 6518–6521. Zhang, C.; Huang, J.; Trudell, M. L.; Nolan, S. P. J. Org. Chem. 1999, 64, 3804–3805. Huang, J.; Nolan, S. P. J. Am. Chem. Soc. 1999, 121, 9889–9890. Yang, C.; Nolan, S. P. Organometallics 2002, 21, 1020–1022. Grasa, G. A.; Viciu, M. S.; Huang, J.; Zhang, C.; Trudell, M. L.; Nolan, S. P. Organometallics 2002, 21, 2866–2873. Viciu, M. S.; Navarro, O.; Germaneau, R. F.; Kelly, R. A., III; Sommer, W.; Marion, N.; Stevens, E. D.; Cavallo, L.; Nolan, S. P. Organometallics 2004, 23, 1629–1635. Altenhoff, G.; Würtz, S.; Glorius, F. Tetrahedron Lett. 2006, 47, 2925–2928. Kantchev, E. A. B.; O’Brien, C. J.; Organ, M. G. Angew. Chem. Int. Ed. 2007, 46, 2768–2813. O’Brien, C. J.; Kantchev, E. A. B.; Valente, C.; Hadei, N.; Chass, G. A.; Lough, A.; Hopkinson, A. C.; Organ, M. G. Chem. Eur. J. 2006, 12, 4743–4748. Organ, M. G.; Abdel-Hadi, M.; Avola, S.; Hadei, N.; Nasielski, J.; O’Brien, C. J.; Valente, C. Chem. Eur. J. 2007, 13, 150–157. Lee, H. M.; Jiang, T.; Stevens, E. D.; Nolan, S. P. Organometallics 2001, 20, 1255–1258. Powell, M. T.; Hou, D.-R.; Perry, M. C.; Cui, X.; Burgess, K. J. Am. Chem. Soc. 2001, 123, 8878. Urban, S.; Ortega, N.; Glorius, F. Angew. Chem. Int. Ed. 2011, 50, 3803–3806. Urban, S.; Beiring, B.; Ortega, N.; Paul, D.; Glorius, F. J. Am. Chem. Soc. 2012, 134, 15241–15244. Ortega, N.; Urban, S.; Beiring, B.; Glorius, F. Angew. Chem. Int. Ed. 2012, 51, 1710–1713. Zhao, D.; Beiring, B.; Glorius, F. Angew. Chem. Int. Ed. 2013, 52, 8454–8458. Wysocki, J.; Ortega, N.; Glorius, F. Angew. Chem. Int. Ed. 2014, 53, 8751–8755. Li, W.; Wiesenfeldt, M. P.; Glorius, F. J. Am. Chem. Soc. 2017, 139, 2585–2588. Dobereiner, G. E.; Nova, A.; Schley, N. D.; Hazari, N.; Miller, S. J.; Eisenstein, O.; Crabtree, R. H. J. Am. Chem. Soc. 2011, 133, 7547–7562. Pranckevicius, C.; Fan, L.; Stephan, D. W. J. Am. Chem. Soc. 2015, 137, 5582–5589. Jochmann, P.; Stephan, D. W. Chem. Eur. J. 2014, 20, 8370–8378. Wei, Y.; Rao, B.; Cong, X.; Zeng, X. J. Am. Chem. Soc. 2015, 137, 9250–9253. Ling, L.; He, Y.; Zhang, X.; Luo, M.; Zeng, X. Angew. Chem. Int. Ed. 2019, 58, 6554–6558. Zhang, X.; Ling, L.; Luo, M.; Zeng, X. Angew. Chem. Int. Ed. 2019, 58, 16785–16789. Wiesenfeldt, M. P.; Nairoukh, Z.; Li, W.; Glorius, F. Science 2017, 357, 908–912. Wiesenfeldt, M. P.; Knecht, T.; Schlepphorst, C.; Glorius, F. Angew. Chem. Int. Ed. 2018, 57, 8297–8300. Wollenburg, M.; Moock, D.; Glorius, F. Angew. Chem. Int. Ed. 2019, 58, 6549–6553. Schneider, S. K.; Herrmann, W. A.; Herdtweck, E. Z. Anorg. Allg. Chem. 2003, 629, 2363–2370. Corpas, J.; Viereck, P.; Chirik, P. J. ACS Catal. 2020, 10, 8640–8647. Marko, I. E.; Sterin, S.; Buisine, O.; Mignani, G.; Branlard, P.; Tinant, B.; Declercq, J.-P. Science 2002, 298, 204–206. Liu, Y.; Deng, L. J. Am. Chem. Soc. 2017, 139, 1798–1801. Ibrahim, A. D.; Entsminger, S. W.; Zhu, L.; Fout, A. R. ACS Catal. 2016, 6, 3589–3593. Liang, Q.; Osten, K. M.; Song, D. Angew. Chem. Int. Ed. 2017, 56, 6317–6320. López, S.; Herrero-Gómez, E.; Pérez-Galán, P.; Nieto-Oberhuber, C.; Echavarren, A. M. Angew. Chem. Int. Ed. 2006, 45, 6029–6032. Escribano-Cuesta, A.; López-Carrillo, V.; Janssen, D.; Echavarren, A. M. Chem. Eur. J. 2009, 15, 5646–5650. Correa, A.; Marion, N.; Fensterbank, L.; Malacria, M.; Nolan, S. P.; Cavallo, L. Angew. Chem. Int. Ed. 2008, 47, 718–721. Marion, N.; de Frémont, P.; Lemière, G.; Stevens, E. D.; Fensterbank, L.; Malacria, M.; Nolan, S. P. Chem. Commun. 2006, 2048–2050. Kim, S. M.; Park, J. H.; Choi, S. Y.; Chung, Y. K. Angew. Chem. Int. Ed. 2007, 46, 6172–6175.

Organometallic Chemistry of NHCs and Analogues 602. 603. 604. 605. 606. 607. 608. 609. 610. 611. 612. 613. 614. 615. 616. 617. 618. 619. 620. 621. 622. 623. 624. 625. 626. 627. 628. 629. 630. 631. 632. 633. 634. 635. 636. 637. 638. 639. 640. 641. 642. 643. 644. 645. 646. 647. 648. 649. 650. 651. 652. 653. 654. 655. 656. 657. 658. 659. 660. 661. 662. 663. 664.

665. 666.

369

Kim, S. M.; Park, J. H.; Kang, Y. K.; Chung, Y. K. Angew. Chem. Int. Ed. 2009, 48, 4532–4535. Witham, C. A.; Mauleo´n, P.; Shapiro, N. D.; Sherry, B. D.; Toste, F. D. J. Am. Chem. Soc. 2007, 129, 5838–5839. Sromek, A. W.; Rubina, M.; Gevorgyan, V. J. Am. Chem. Soc. 2005, 127, 10500–10501. Dudnik, A. S.; Gevorgyan, V. Angew. Chem. Int. Ed. 2007, 46, 5195–5197. Zhou, C.-Y.; Chan, P. W. H.; Che, C. M. Org. Lett. 2006, 8, 325–328. Zhang, G.; Huang, X.; Li, G.; Zhang, L. J. Am. Chem. Soc. 2008, 130, 1814–1815. Li, G.; Huang, X.; Zhang, L. J. Am. Chem. Soc. 2008, 130, 6944–6945. Peng, Y.; Yu, M.; Zhang, L. Org. Lett. 2008, 10, 5187–5190. Lavallo, V.; Frey, G. D.; Kousar, S.; Donnadieu, B.; Bertrand, G. Proc. Natl. Acad. Sci. U.S.A. 2007, 104, 13569–13573. Lavallo, V.; Frey, G. D.; Kousar, S.; Donnadieu, B.; Soleilhavoup, M.; Bertrand, G. Angew. Chem. Int. Ed. 2008, 47, 5224–5228. Zeng, X.; Frey, G. D.; Kinjo, R.; Donnadieu, B.; Bertrand, G. J. Am. Chem. Soc. 2009, 131, 8690–8696. Kaufhold, S.; Petermann, L.; Staehle, R.; Rau, S. Coord. Chem. Rev. 2015, 304–305, 73–87. Ohki, Y.; Seinoc, H. Dalton Trans. 2016, 45, 874–880. Zhang, L.; Li, Z.; Takimoto, M.; Hou, Z. Chem. Rec. 2020, 20, 494–512. Albrecht, M.; Miecznikowski, J. R.; Samuel, A.; Faller, J. W.; Crabtree, R. H. Organometallics 2002, 21, 3596–3604. Su, X.; McCardle, K. M.; Panetier, J. A.; Jurss, J. W. Chem. Commun. 2018, 54, 3351–3354. Peuntinger, K.; David Pilz, T.; Staehle, R.; Schaub, M.; Kaufhold, S.; Petermann, L.; Wunderlin, M.; Görls, H.; Heinemann, F. W.; Li, J.; Drewello, T.; Vos, J. G.; Guldi, D. M.; Rau, S. Dalton Trans. 2014, 43, 13683–13695. Codolà, Z.; Cardoso, J. M. S.; Royo, B.; Costas, M.; Lloret-Fillol, J. Chem. Eur. J. 2013, 19, 7203–7213. Lalrempuia, R.; McDaniel, N. D.; Müller-Bunz, H.; Bernhard, S.; Albrecht, M. Angew. Chem. Int. Ed. 2010, 49, 9765–9768. Ashida, Y.; Arashiba, K.; Nakajima, K.; Nishibayashi, Y. Nature 2019, 568. 563-540. Mercsa, L.; Albrecht, M. Chem. Soc. Rev. 2010, 39, 1903–1912. Pernak, J.; Skrzypczak, A. Eur. J. Med. Chem. 1996, 31, 901–903. Aher, S. B.; Muskawar, P. N.; Thenmozhi, K.; Bhagat, P. R. Eur. J. Med. Chem. 2014, 81, 408–419. Gautier, A.; Cisnetti, F. Metallomics 2012, 4, 23–32. Teyssot, M. L.; Jarrousse, A. S.; Manin, M.; Chevry, A.; Roche, S.; Norre, F.; Beaudoin, C.; Morel, L.; Boyer, D.; Mahiou, R.; Gautier, A. Dalton Trans. 2009, 6894–6902. Johnson, N. A.; Southerl, M. R.; Youngs, W. J. Molecules 2017, 22, 1263–1282. Karaaslan, M. G.; Aktas¸, A.; Gürses, C.; Gök, Y.; Ate, B. Bioorg. Chem. 2020, 95, 103552–103565. Sivaram, H.; Tan, J.; Huynh, H. V. Organometallics 2012, 31, 5875–5883. Kumar, A.; Naaz, A.; Prakasham, A. P.; Gangwar, M. K.; Butcher, R. J.; Panda, D.; Ghosh, P. ACS Omega 2017, 2, 4632–4646. Alfaro, J. M.; Prades, A.; Ramos, M. C.; Peris, E.; Repoll-Grómes, J.; Poyatos, M.; Burgos, J. S. Zebrafish 2010, 7, 13–21. Melaiye, A.; Simons, R. S.; Milsted, A.; PingItore, F.; Wesdemiotis, C.; Tessier, C. A.; Youngs, W. J. J. Med. Chem. 2004, 47, 973–977. Kaps, L.; Biersack, B.; Müller-bunz, H.; Mahal, K.; Münzner, J.; Tacke, M.; Mueller, T.; Schobert, R. J. Inorg. Biochem. 2012, 106, 52–58. Kascatan-Nebioglu, A.; Melaiye, A.; Hindi, K.; Durmus, S.; Panzner, M. J.; Hogue, L. A.; Mallet, R. J.; Hovis, C. E.; Coughenous, M.; Crosby, S. D.; Milsted, A.; Ely, D. L.; Tessier, C. A.; Cannon, C. L.; Youngs, W. J. J. Med. Chem. 2006, 49, 6811–6818. Mohamed, H. A.; Shepherd, S.; William, N.; Blundell, H. A.; Das, M.; Pask, C. M.; Lake, B. M.; Phillips, R. M.; Nelson, A.; Willans, C. E. Organometallics 2020, 39, 1318–1331. Lemke, J.; Pinto, A.; Niehoff, P.; Vasylyeva, V.; Meltzler-Nolte, N. Dalton Trans. 2009, 7063–7070. Sun, R. W.-Y.; Chow, A. L.-F.; Li, X.-H.; Yan, J. J.; Chui, S. S.-Y.; Che, C.-M. Chem. Sci. 2011, 2, 728–736. Tsui, W.-K.; Chung, L.-H.; Wong, M. M.-K.; Tsang, W.-H.; Lo, H.-S.; Liu, Y.-X.; Leung, C.-H.; Ma, D.-L.; Chiu, S. K.; Wong, C.-Y. Sci. Rep. 2015, 5, 9070–9074. Yam, V. W. W.; Wong, K. M. C. Chem. Commun. 2011, 47, 11579–11592. Elie, M.; Renaud, J. L.; Gaillard, S. Polyhedron 2018, 140, 158–168. Chábera, P.; Lindh, L.; Rosemann, N. W.; Prakash, O.; Uhlig, J.; Yartsev, A.; Wärnmark, K.; Sundström, V.; Persson, P. Coord. Chem. Rev. 2021, 426, 213517. Wenger, O. S. J. Am. Chem. Soc. 2018, 140, 13522–13533. Förster, C.; Heinze, K. Chem. Soc. Rev. 2020, 49, 1057–1070. Xue, W. M.; Chan, M. C. W.; Su, Z. M.; Cheung, K. K.; Liu, S. T.; Che, C. M. Organometallics 1998, 17, 1622–1630. Son, S. U.; Park, K. H.; Lee, Y. S.; Kim, B. Y.; Choi, C. H.; Lah, M. S.; Jang, Y. H.; Jang, D. J.; Chung, Y. K. Inorg. Chem. 2004, 43, 6896–6898. Duan, G. P.; Yam, V. W. W. Chem. Eur. J. 2010, 16, 12642–12649. Schulze, B.; Escudero, D.; Friebe, C.; Siebert, R.; Gçrls, H.; Kçhn, U.; Altuntas, E.; Baumgaertel, A.; Hager, M. D.; Winter, A.; Benjamin Dietzek, B.; Popp, J.; Leticia Gonzalez, L.; Schubert, U. S. Chem. Eur. J. 2011, 17, 5494–5498. Holmes, R. J.; Forrest, S. R.; Sajoto, T.; Tamayo, A.; Djurovich, P. I.; Thompson, M. E.; Brooks, J.; Tung, Y.-J.; D’Andrade, B. W.; Weaver, M. S.; Kwong, R. C.; Brown, J. J. Appl. Phys. Lett. 2005, 87, 243507. Darmawan, N.; Yang, C. H.; Mauro, M.; Raynal, M.; Heun, S.; Pan, J. Y.; Buchholz, H.; Braunstein, P.; Cola, L. D. Inorg. Chem. 2013, 52, 10756–10765. Lu, K. Y.; Chou, H. H.; Hsieh, C. H.; Ou Yang, Y. H.; Tsai, H. R.; Tsai, H. Y.; Hsu, L. C.; Chen, C. Y.; Chen, I. C.; Cheng, C. H. Adv. Mater. 2011, 23, 4933–4937. Lee, J.; Chen, H. F.; Batagoda, T.; Coburn, C.; Djurovich, P. I.; Thompson, M. E.; Forrest, S. R. Nature. Mater. 2016, 15, 92–99. Li, K.; Chen, Y.; Lu, W.; Zhu, N. Y.; Che, C. M. Chem. Eur. J. 2011, 17, 4109–4112. Lee, C. S.; Zhuang, R. R.; Sabiah, S.; Wang, J. C.; Hwang, W. S.; Lin, I. J. B. Organometallics 2011, 30, 3897–3900. Leung, S. Y. L.; Lam, E. S. H.; Lam, W. H.; Wong, K. M. C.; Wong, W. T.; Yam, V. W. W. Chem. Eur. J. 2013, 19, 10360–10369. Au, V. K. M.; Wong, K. M. C.; Zhu, N. Y.; Yam, V. W. W. J. Am. Chem. Soc. 2009, 131, 9076–9085. Yam, V. W. W.; Lee, J. K. W.; Ko, C. C.; Zhu, N. Y. J. Am. Chem. Soc. 2009, 131, 912–913. Unger, Y.; Zeller, A.; Ahrensy, S.; Strassner, T. Chem. Commun. 2008, 3263–3265. Visbal, R.; Ospino, I.; López-de-Luzuriaga, J. M.; Laguna, A.; Gimeno, M. C. J. Am. Chem. Soc. 2013, 135, 4712–4715. Catalano, V. J.; Moore, A. L. Inorg. Chem. 2005, 44, 6558–6566. Gimeno, M. C.; Laguna, A.; Visbal, R. Organometallics 2012, 31, 7146–7157. Matsumoto, K.; Matsumoto, N.; Ishii, A.; Tsukuda, T.; Hasegawa, M.; Tsubomura, T. Dalton Trans. 2009, 6795–6801. Nitsch, J.; Lacemon, F.; Lorbach, A.; Eichhorn, A.; Cisnetti, F.; Steffen, A. Chem. Commun. 2016, 52, 2932–2935. Hamze, R.; Peltier, J. L.; Sylvinson, D.; Jung, M.; Cardenas, J.; Haiges, R.; Soleilhavoup, M.; Jazzar, R.; Djurovich, P. I.; Bertrand, G.; Thompson, M. E. Science 2019, 363, 601–606. Chábera, P.; Liu, Y. Z.; Prakash, O.; Thyrhaug, E.; Nahhas, A. E.; Honarfar, A.; Essén, S.; Fredin, L. A.; Harlang, T. C. B.; Kjær, K. S.; Handrup, K.; Ericson, F.; Tatsuno, H.; Morgan, K.; Schnadt, J.; Häggström, L.; Ericsson, T.; Sobkowiak, A.; Lidin, S.; Huang, P.; Styring, S.; Uhlig, J.; Bendix, J.; Lomoth, R.; Sundström, V.; Persson, P.; Wärnmark, K. Nature 2017, 543, 695–699. Kjær, K. S.; Kaul, N.; Prakash, O.; Chábera, P.; Rosemann, N. W.; Honarfar, A.; Gordivska, O.; Fredin, L. A.; Bergquist, K. E.; Häggström, L.; Ericsson, T.; Lindh, L.; Yartsev, A.; Styring, S.; Huang, P.; Uhlig, J.; Bendix, J.; Strand, D.; Sundström, V.; Persson, P.; Lomoth, R.; Wärnmark, K. Science 2019, 363, 249–253. Altmann, P. J.; Pöthig, A. J. Am. Chem. Soc. 2016, 138, 13171–13174.

370 667. 668. 669. 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. 696. 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. 723. 724. 725. 726. 727. 728. 729. 730. 731. 732. 733. 734.

Organometallic Chemistry of NHCs and Analogues Altmann, P. J.; Pöthig, A. Angew. Chem. Int. Ed. 2017, 56, 15733–15736. Nishad, R. C.; Rit, A. Chem. Eur. J. https://doi.org/10.1002/chem.202003z937. Pöthig, A.; Casini, A. Theranostics 2019, 9, 3150–3169. Modak, R.; Mondal, B.; Howlader, P.; Mukherjee, P. S. Chem. Commun. 2019, 55, 6711–6714. Wang, D.; Zhang, B.; He, C.; Wu, P.; Duan, C. Chem. Commun. 2010, 46, 4728–4730. Zhukhovitskiy, A. V.; MacLeod, M. J.; Johnson, J. A. Chem. Rev. 2015, 115, 11503–11532. Zhukhovitskiy, A. V.; Mavros, M. G.; Van Voorhis, T.; Johnson, J. A. J. Am. Chem. Soc. 2013, 135, 7418–7421. Crudden, C. M.; Horton, J. H.; Ebralidze, I. I.; Zenkina, O. V.; McLean, A. B.; Drevniok, B.; She, Z.; Kraatz, H. B.; Mosey, N. J.; Seki, T.; Keske, E. C.; Leake, J. D.; RousinaWebb, A.; Wu, G. Nat. Chem. 2014, 6, 409–414. Ranganath, K. V. S.; Kloesges, J.; Schäfer, A. H.; Glorius, F. Angew. Chem. Int. Ed. 2010, 49, 7786–7789. Ernst, J. B.; Muratsugu, S.; Wang, F.; Tada, M.; Glorius, F. J. Am. Chem. Soc. 2016, 138, 10718–10721. Wang, G.; Rühling, A.; Amirjalayer, S.; Knor, M.; Ernst, J. B.; Richter, C.; Gao, H.-J.; Timmer, A.; Gao, H.-Y.; Doltsinis, N. L.; Glorius, F.; Fuchs, H. Nat. Chem. 2017, 9, 152–156. Amirjalayer, S.; Bakker, A.; Freitag, M.; Glorius, F.; Fuchs, H. Angew. Chem. Int. Ed. 2020, 59, 21230–21235. Oisaki, K.; Li, Q.; Furukawa, H.; Czaja, A. U.; Yaghi, O. M. J. Am. Chem. Soc. 2010, 132, 9262–9264. Kong, G.-Q.; Ou, S.; Zou, C.; Wu, C.-D. J. Am. Chem. Soc. 2012, 134, 19851–19857. Burgun, A.; Crees, R. S.; Cole, M. L.; Doonan, C. J.; Sumby, C. J. Chem. Commun. 2014, 50, 11760–11763. Crees, R. S.; Cole, M. L.; Hanton, L. R.; Sumby, C. J. Inorg. Chem. 2010, 49, 1712–1719. Wang, Y.; Robinson, G. H. Inorg. Chem. 2014, 53, 11815–11832. Wang, Y.; Robinson, G. H. Inorg. Chem. 2011, 50, 12326–12337. Kundu, S.; Sinhababu, S.; Chandrasekhar, V.; Roesky, H. W. Chem. Sci. 2019, 10, 4727–4741. Zhao, L.; Pan, S.; Holzmann, N.; Schwerdtfeger, P.; Frenking, G. Chem. Rev. 2019, 119, 8781–8845. Wang, Y.; Xie, Y.; Wei, P.; King, R. B.; Schaefer, H. F., III; Schleyer, P.; von., R.; Robinson, G. H. Science 2008, 321, 1069–1071. Sidiropoulos, A.; Jones, C.; Stasch, A.; Klein, S.; Frenking, G. Angew. Chem. Int. Ed. 2009, 48, 9701–9704. Jones, C.; Sidiropoulos, A.; Holzmann, N.; Frenking, G.; Stasch, A. Chem. Commun. 2012, 48, 9855–9857. Braunschweig, H.; Dewhurst, R. D.; Hammond, K.; Mies, J.; Radacki, K.; Vargas, A. Science 2012, 336, 1420–1422. Dyker, C.; Lavallo, V.; Donnadieu, B.; Bertrand, G. Angew. Chem. Int. Ed. 2008, 47, 3206–3209. Mondal, K. C.; Roesky, H. W.; Schwarzer, M. C.; Frenking, G.; Niepötter, B.; Wolf, H.; Herbst-Irmer, R.; Stalke, D. Angew. Chem. Int. Ed. 2013, 52, 2963–2967. Li, Y.; Mondal, K. C.; Roesky, H.; Zhu, P.; Stollberg, P.; HerbstIrmer, R.; Stalke, D.; Andrada, D. M. J. Am. Chem. Soc. 2013, 135, 12422–12428. Xiong, Y.; Yao, S.; Inoue, S.; Epping, J. D.; Driess, M. Angew. Chem. Int. Ed. 2013, 52, 7147–7150. Xiong, Y.; Yao, S.; Tan, G.; Inoue, S.; Driess, M. J. Am. Chem. Soc. 2013, 135, 5004–5007. Su, B.; Ganguly, R.; Li, Y.; Kinjo, R. Angew. Chem. Int. Ed. 2014, 53, 13106–13109. Kinjo, R.; Donnadieu, B.; Ali Celik, M.; Frenking, G.; Bertrand, G. Science 2011, 333, 610–613. Légaré, M.-A.; Bélanger-Chabot, G.; Rang, M.; Dewhurst, R. D.; Krummenacher, I.; Bertermann, R.; Braunschweig, H. Nat. Chem. 2020, 12, 1076–1080. Légaré, M.-A.; Bélanger-Chabot, G.; Dewhurst, R. D.; Welz, E.; Krummenacher, I.; Engels, B.; Braunschweig, H. Science 2018, 359, 896–900. Légaré, M.-A.; Rang, M.; Bélanger-Chabot, G.; Schweizer, J. I.; Krummenacher, I.; Bertermann, R.; Arrowsmith, M.; Holthausen, M. C.; Braunschweig, H. Science 2019, 363, 1329–1332. Crudden, C. M.; Allen, D. P. Coord. Chem. Rev. 2004, 248, 2247–2273. Würtemberger-Pietsch, S.; Radius, U.; Marder, T. B. Dalton Trans. 2016, 45, 5880–5895. Chernyshev, V. M.; Denisova, E. A.; Eremin, D. B.; Ananikov, V. P. Chem. Sci. 2020, 11, 6957–6977. Hu, X.; Meyer, K. J. Am. Chem. Soc. 2004, 126, 16322–16323. Zolnhofer, E. M.; Käß, M.; Khusniyarov, M. M.; Heinemann, F. W.; Maron, L.; van Gastel, M.; Bill, E.; Meyer, K. J. Am. Chem. Soc. 2014, 136, 15072–15078. Liu, B.; Zhang, Y.; Xu, D.; Chen, W. Chem. Commun. 2011, 47, 2883–2885. Danopoulos, A. A.; Tsoureas, N.; Green, J. C.; Hursthouse, M. B. Chem. Commun. 2003, 756–757. McGuinness, D. S.; Green, M. J.; Cavell, K. J.; Skelton, B. W.; White, A. H. J. Organomet. Chem. 1998, 565, 165. McGuinness, D. S.; Cavell, K. J.; Skelton, B. W.; White, A. H. Organometallics 1999, 18, 1596–1605. Lin, B.-L.; Kang, P.; Stack, T. D. P. Organometallics 2010, 29, 3683–3685. Khazipov, O. V.; Shevchenko, M. A.; Chernenko, A. Y.; Astakhov, A. V.; Pasyukov, D. V.; Eremin, D. B.; Zubavichus, Y. V.; Khrustalev, V. N.; Chernyshev, V. M.; Ananikov, V. P. Organometallics 2018, 37, 1483–1492. Wegner, S.; Janiak, C. Top. Curr. Chem. 2017, 375, 65. Eremin, D. B.; Denisova, E. A.; Kostyukovich, A. Y.; Martens, J.; Berden, G.; Oomens, J.; Khrustalev, V. N.; Chernyshev, V. M.; Ananikov, V. P. Chem. Eur. J. 2019, 25, 16564–16572. Tran, B. L.; Fulton, J. L.; Linehan, J. C.; Balasubramanian, M.; Lercher, J. A.; Bullock, R. M. ACS Catal. 2019, 9, 4106–4114. Tran, B. L.; Fulton, J. L.; Linehan, J. C.; Lercher, J. A.; Bullock, R. M. ACS Catal. 2018, 8, 8441–8449. Asensio, J. M.; Tricard, S.; Coppel, Y.; Andŕes, R.; Chaudret, B.; de Jesus, E. Chem. Eur. J. 2017, 23, 13435–13444. Huang, J.; Stevens, E. D.; Nolan, S. P. Organometallics 2000, 19, 1194–1197. Sun, J.; Ou, C.; Wang, C.; Uchiyama, M.; Deng, L. Organometallics 2015, 34, 1546–1551. Torres, O.; Martín, M.; Sola, E. Organometallics 2009, 28, 863–870. Choi, G.; Tsurugi, H.; Mashima, K. J. Am. Chem. Soc. 2002, 135, 13149–13161. Ohki, Y.; Hatanaka, T.; Tatsumi, K. J. Am. Chem. Soc. 2008, 130, 17174–17186. Rivada-Wheelaghan, O.; Ortunño, M. A.; Dίez, J.; Lledόs, A.; Conejero, S. Angew. Chem. Int. Ed. 2012, 51, 3936–3939. Tang, C. Y.; Smith, W.; Thompson, A. L.; Vidovic, D.; Aldridge, S. Angew. Chem. Int. Ed. 2011, 50, 1359–1362. Navarro, J.; Torres, O.; Martín, M.; Sola, E. J. Am. Chem. Soc. 2011, 133, 9738–9740. Rivada-Wheelaghan, O.; Ortunño, M. A.; Dίez, J.; Garcia-Garrido, S. E.; Maya, C.; Lledόs, A.; Conejero, S. J. Am.Chem. Soc. 2012, 134, 15261–15264. Mo, Z.; Chen, D.; Leng, X.; Deng, L. Organometallics 2012, 31, 7040–7043. Ouyang, Z.; Deng, L. Organometallics 2013, 32, 7268–7271. Gao, Y.; Chen, Q.; Leng, X.; Deng, L. Dalton Trans. 2019, 48, 9676–9683. Jazzar, R. F. R.; Macgregor, S. A.; Mahon, M. F.; Richards, S. P.; Whittlesey, M. K. J. Am. Chem. Soc. 2002, 124, 4944–4945. Bolano, T.; Buil, M. L.; Esteruelas, M. A.; Izquierdo, S.; Lalrempuia, R.; Olivan, M.; Onate, E. Organometallics 2010, 29, 4517–4523. Häller, L. J. L.; Page, M. J.; Erhardt, S.; Macgregor, S. A.; Mahon, M. F.; Naser, M. A.; Vélez, A.; Whittlesey, M. K. J. Am. Chem. Soc. 2010, 132, 18408–18416. Burling, S.; Mahon, M. F.; Powell, R. E.; Whittlesey, M. K.; Williams, J. M. J. J. Am. Chem. Soc. 2006, 128, 13702–13703. Caddick, S.; Cloke, F. G. N.; Hitchcock, P. B.; de, K.; Lewis, A. K. Angew. Chem. Int. Ed. 2004, 43, 5824–5827. Xiao, J.; Deng, L. Organometallics 2012, 31, 428–434.

Organometallic Chemistry of NHCs and Analogues 735. 736. 737. 738. 739. 740. 741. 742. 743. 744. 745. 746. 747. 748. 749. 750. 751. 752. 753. 754. 755. 756. 757. 758. 759. 760. 761. 762. 763. 764. 765. 766. 767. 768. 769. 770. 771. 772. 773. 774. 775. 776. 777. 778. 779. 780. 781. 782. 783. 784. 785. 786. 787. 788. 789. 790. 791. 792. 793. 794. 795. 796. 797. 798. 799. 800. 801. 802. 803. 804. 805. 806.

371

Day, B. M.; Pugh, T.; Hendriks, D.; Guerra, C. F.; Evans, D. J.; Bickelhaupt, F. M.; Layfield, R. A. J. Am. Chem. Soc. 2013, 135, 13338–13341. Hering, F.; Radius, U. Organometallics 2015, 34, 3236–3245. Sakurai, S.; Tobisu, M. Organometallics 2019, 38, 2834–2838. Schmidt, D.; Berthel, J. H. J.; Pietsch, S.; Radius, U. Angew. Chem. Int. Ed. 2012, 51, 8881–8885. Arrowsmith, M.; Hill, M. S.; Kociok-Köhn, G.; MacDougall, D. J.; Mahon, M. F. Angew. Chem. Int. Ed. 2012, 51, 2098–2100. Al-Rafia, S. M. I.; McDonald, R.; Ferguson, M. J.; Rivard, E. Chem. – Eur. J. 2012, 18, 13810–13820. Franz, D.; Inoue, S. Chem. – Asian J. 2014, 9, 2083–2087. Wang, T.; Stephan, D. W. Chem. – Eur. J. 2014, 20, 3036–3039. Pelegri, A. S.; Elsegood, M. R. J.; McKee, C.; Weaver, G. W. Org. Lett. 2006, 8, 3049–3051. Waltman, A. W.; Ritter, T.; Grubbs, R. H. Organometallics 2006, 25, 4238–4239. Liu, H.-J.; Ziegler, M. S.; Tilley, T. D. Polyhedron 2014, 84, 203–208. Danopoulos, A. A. Dalton Trans. 2008, 1087–1094. Hatanaka, T.; Ohki, Y.; Tatsumi, K. Angew. Chem. Int. Ed. 2014, 53, 2727–2729. Zuo, W.; Braunstein, P. Dalton Trans. 2012, 41, 636–643. Prema, D.; Mathota Arachchige, Y. L. N.; Murray, R. E.; Slaughter, L. M. Chem. Commun. 2015, 51, 6753–6756. Asay, M.; Jones, C.; Driess, M. Chem. Rev. 2011, 111, 354–396. Segawa, Y.; Yamashita, M.; Nozaki, K. Science 2006, 314, 113–115. Segawa, Y.; Suzuki, Y.; Yamashita, M.; Nozaki, K. J. Am. Chem. Soc. 2008, 130, 16069–16079. Yamashita, M.; Suzuki, Y.; Segawa, Y.; Nozaki, K. Chem. Lett. 2008, 37, 802–803. Lu, W.; Hu, H.; Li, Y.; Ganguly, R.; Kinjo, R. J. Am. Chem. Soc. 2016, 138, 6650–6661. Segawa, Y.; Yamashita, M.; Nozaki, K. Angew. Chem. Int. Ed. 2007, 46, 6710–6713. Kajiwara, T.; Terabayashi, T.; Yamashita, M.; Nozaki, K. Angew. Chem. Int. Ed. 2008, 47, 6606–6610. Terabayashi, T.; Kajiwara, T.; Yamashita, M.; Nozaki, K. J. Am. Chem. Soc. 2009, 131, 14162–14163. Saleh, L. M. A.; Birjkumar, K. H.; Protchenko, A. V.; Schwarz, A. D.; Aldridge, S.; Jones, C.; Kaltsoyannis, N.; Mountford, P. J. Am. Chem. Soc. 2011, 133, 3836–3839. Li, S.; Cheng, J.; Chen, Y.; Nishiura, M.; Hou, Z. Angew. Chem. Int. Ed. 2011, 50, 6360–6363. Okuno, Y.; Yamashita, M.; Nozaki, K. Angew. Chem. Int. Ed. 2011, 50, 920–923. Ga, In, Tl: Protchenko, A. V.; Dange, D.; Harmer, J. R.; Tang, C. Y.; Schwarz, A. D.; Kelly, M. J.; Phillips, N.; Tirfoin, R.; Birjkumar, K. H.; Jones, C.; Kaltsoyannis, N.; Mountford, P.; Aldridge, S. Nat. Chem. 2014, 6, 315 − 319. Protchenko, A. V.; Birjkumar, K. H.; Dange, D.; Schwarz, A. D.; Vidovic, D.; Jones, C.; Kaltsoyannis, N.; Mountford, P.; Aldridge, S. J. Am. Chem. Soc. 2012, 134, 6500–6503. Rit, A.; Campos, J.; Niu, H.; Aldridge, S. Nat. Chem. 2016, 8, 1022–1026. Cui, C.; Roesky, H. W.; Schmidt, H.-G.; Noltemeyer, M.; Hao, H.; Cimpoesu, F. Angew. Chem. Int. Ed. 2000, 39, 4274–4276. Li, X.; Cheng, X.; Song, H.; Cui, C. Organometallics 2007, 26, 1039–1043. Zhu, H.; Chai, J.; Jancik, V.; Roesky, H. W.; Merrill, W. A.; Power, P. P. J. Am. Chem. Soc. 2005, 127, 10170–10171. Chu, T.; Korobkov, I.; Nikonov, G. I. J. Am. Chem. Soc. 2014, 136, 9195–9202. Bakewell, C.; White, A. J. P.; Crimmin, M. R. Angew. Chem. Int. Ed. 2018, 57, 6638–6642. Bakewell, C.; White, A. J. P.; Crimmin, M. R. Chem. Sci. 2019, 10, 2452–2458. Kong, R. Y.; Crimmin, M. R. J. Am. Chem. Soc. 2018, 140, 13614–13617. Jones, C.; Junk, P. C.; Platts, J. A.; Stasch, A. J. Am. Chem. Soc. 2006, 128, 2206–2207. Hardman, N. J.; Eichler, B. E.; Power, P. P. Chem. Commun. 2000, 1991–1992. Stender, M.; Power, P. P. Polyhedron 2002, 21, 525–529. Hill, M. S.; Pontavornpinyo, R.; Hitchcock, P. B. Chem. Commun. 2006, 3720–3722. Hicks, J.; Vasko, P.; Goicoechea, J. M.; Aldridge, S. Nature 2018, 557, 92–95. Hicks, J.; Mansikkamaki, A.; Vasko, P.; Goicoechea, J. M.; Aldridge, S. Nat. Chem. 2019, 11, 237–241. Schwamm, R. J.; Anker, M. D.; Lein, M.; Coles, M. P. Angew. Chem. Int. Ed. 2019, 58, 1489–1493. Schwamm, R. J.; Coles, M. P.; Hill, M. S.; Mahon, M. F.; McMullin, C. L.; Rajabi, N. A.; Wilson, A. S. S. Angew. Chem. Int. Ed. 2020, 59, 3928–3932. Kurumada, S.; Takamori, S.; Yamashita, M. Nat. Chem. 2020, 12, 36–39. Koshino, K.; Kinjo, R. J. Am. Chem. Soc. 2020, 142, 9057–9062. Schmidt, E. S.; Jockisch, A.; Schmidbaur, H. J. Am. Chem. Soc. 1999, 121, 9758–9759. Baker, R. J.; Jones, C.; Mills, D. P.; Pierce, G. A.; Waugh, M. Inorg. Chim. Acta 2008, 361, 427–435. Fedushkin, I.; Lukoyanov, A. N.; Fukin, G. K.; Ketkov, S. Y.; Hummert, M.; Schumann, H. Chem.-Eur. J. 2008, 14, 8465–8468. Fedushkin, I. L.; Lukoyanov, A. N.; Tishkina, A. N.; Fukin, G. K.; Lyssenko, K. A.; Hummert, M. Chem.-Eur. J. 2010, 16, 7563–7566. Schwamm, R. J.; Anker, M. D.; Lein, M.; Coles, M. P.; Fitchett, C. M. Angew. Chem. Int. Ed. 2018, 57, 5885–5887. Zhou, Y.-P.; Driess, M. Angew. Chem. Int. Ed. 2019, 58, 3715–3728. Denk, M.; Lennon, R.; Hayashi, R.; West, R.; Belyakov, A. V.; Verne, H. P.; Haaland, A.; Wagner, M.; Metzler, N. J. Am. Chem. Soc. 1994, 116, 2691–2692. West, R.; Denk, M. Pure Appl. Chem. 1996, 68, 785–788. Schmedake, T. A.; Haaf, M.; Apeloig, Y.; Mueller, T.; Bukalov, S.; West, R. J. Am. Chem. Soc. 1999, 121, 9479–9480. Li, W.; Hill, N. J.; Tomasik, A. C.; Bikzhanova, G.; West, R. Organometallics 2006, 25, 3802–3805. Tomasik, A. C.; Mitra, A.; West, R. Organometallics 2009, 28, 378–381. Kong, L.; Zhang, J.; Song, H.; Cui, C. Dalton Trans. 2009, 5444–5446. Driess, M.; Yao, S.; Brym, M.; van Wuellen, C.; Lentz, D. J. Am. Chem. Soc. 2006, 128, 9628–9629. Kosai, T.; Ishida, S.; Iwamoto, T. Angew. Chem. Int. Ed. 2016, 55, 15554–15558. So, C.-W.; Roesky, H. W.; Magull, J.; Oswald, R. B. Angew. Chem. Int. Ed. 2006, 45, 3948–3950. Wang, W.; Inoue, S.; Yao, S.; Driess, M. J. Am. Chem. Soc. 2010, 132, 15890–15892. Wang, W.; Inoue, S.; Enthaler, S.; Driess, M. Angew. Chem. Int. Ed. 2012, 51, 6167–6171. Zhou, Y.-P.; Raoufmoghaddam, S.; Szilvási, T.; Driess, M. Angew. Chem. Int. Ed. 2016, 55, 12868–12872. Wang, W.; Inoue, S.; Irran, E.; Driess, M. Angew. Chem. Int. Ed. 2012, 51, 3691–3694. Brück, A.; Gallego, D.; Wang, W.; Irran, E.; Driess, M.; Hartwig, J. F. Angew. Chem. Int. Ed. 2012, 51, 11478–11482. Gallego, D.; Brück, A.; Irran, E.; Meier, F.; Kaupp, M.; Driess, M.; Hartwig, J. F. J. Am. Chem. Soc. 2013, 135, 15617–15626. Gallego, D.; Inoue, S.; Blom, B.; Driess, M. Organometallics 2014, 33, 6885–6897. Wang, W.; Kostenko, A.; Yao, S.; Driess, M. J. Am. Chem. Soc. 2017, 139, 13499–13506. Zhong, F.; Yang, X.; Shen, L.; Zhao, Y.; Ma, H.; Wu, B.; Yang, X.-J. Inorg. Chem. 2016, 55, 9112–9120. Benedek, Z.; Szilvási, T. RSC Adv. 2015, 5, 5077–5086. Meltzer, A.; Inoue, S.; Pr-sang, C.; Driess, M. J. Am. Chem. Soc. 2010, 132, 3038–3046.

372 807. 808. 809. 810. 811. 812. 813. 814. 815. 816. 817. 818. 819. 820. 821. 822. 823. 824. 825. 826. 827. 828.

Organometallic Chemistry of NHCs and Analogues Fürstner, A.; Krause, H.; Lehmann, C. W. Chem. Commun. 2001, 2372–2373. Denk, M. K.; Gupta, S.; Ramachandran, R. Tetrahedron Lett. 1996, 37, 9025–9028. Carmalt, C. J.; Lomeli, V.; McBurnett, B. G.; Cowley, A. H. Chem. Commun. 1997, 2095–2096. Gudat, D.; Gans-Eichler, T.; Nieger, M. Chem. Commun. 2004, 2434–2435. Tulchinsky, Y.; Iron, M. A.; Botoshansky, M.; Gandelman, M. Nat. Chem. 2011, 3, 525–531. Day, G. S.; Pan, B.; Kellenberger, D. L.; Foxman, B. M.; Thomas, C. M. Chem. Commun. 2011, 47, 3634–3636. Abrams, M. B.; Scott, B. L.; Baker, R. T. Organometallics 2000, 19, 4944–4956. Spinney, H. A.; Yap, G. P. A.; Korobkov, I.; DiLabio, G.; Richeson, D. S. Organometallics 2006, 25, 3541–3543. Gudat, D. Coord. Chem. Rev. 1997, 163, 71–106. Nakazawa, H. Adv. Organomet. Chem. 2004, 50, 107–143. Rosenberg, L. Coord. Chem. Rev. 2012, 256, 606–626. Gudat, D. Low-Coordinate Main Group Compounds-Group 15. In Comprehensive Inorganic Chemistry II; Reedijk, J.; Poeppelmeier, K., Eds.; Elsevier: Oxford, 2013, Vol 1, pp 587–621. Hardman, N. J.; Abrams, M. B.; Pribisko, M. A.; Gilbert, T. M.; Martin, R. L.; Kubas, G. J.; Baker, R. T. Angew. Chem. Int. Ed. 2004, 43, 1955–1958. Nakazawa, H.; Miyoshi, Y.; Katayama, T.; Mizuta, T.; Miyoshi, K.; Tsuchida, N.; Ono, A.; Takano, K. Organometallics 2006, 25, 5913–5921. Caputo, C. A.; Brazeau, A. L.; Hynes, Z.; Price, J. T.; Tuononen, H. M.; Jones, N. D. Organometallics 2009, 28, 5261–5265. Caputo, C. A.; Jennings, M. C.; Tuononen, H. M.; Jones, N. D. Organometallics 2009, 28, 990–1000. Price, J. T.; Lui, M.; Jones, N. D.; Ragogna, P. J. Inorg. Chem. 2011, 50, 12810–12817. Pan, B.; Bezpalko, M. W.; Foxman, B. M.; Thomas, C. M. Organometallics 2011, 30, 5560–5563. Förster, D.; Nickolaus, J.; Nieger, M.; Benko˝ , Z.; Ehlers, A. W.; Gudat, D. Inorg. Chem. 2013, 52, 7699–7708. Pan, B.; Xu, Z.; Bezpalko, M. W.; Foxman, B. M.; Thomas, C. M. Inorg. Chem. 2012, 51, 4170–4179. Pan, B.; Bezpalko, M. W.; Foxman, B. M.; Thomas, C. M. Dalton Trans. 2012, 41, 9083–9090. Tulchinsky, Y.; Kozuch, S.; Saha, P.; Mauda, A.; Nisnevich, G.; Botoshansky, M.; Shimon, L. J. W.; Gandelman, M. Chem. Eur. J. 2015, 21, 7099–7110.

1.12 Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions Anton W Tomicha, Varun Teja, Sergio Loveraa, Isaac Bandaa, Steven Fisherb, Matthew Asayc, and Vincent Lavalloa, aUniversity of California Riverside, Riverside, CA, United States; bLawrence Livermore National Laboratory, Livermore, CA, United States; cUniversal Display Corporation, Ewing, NJ, United States © 2022 Elsevier Ltd. All rights reserved.

1.12.1 1.12.2 1.12.2.1 1.12.2.1.1 1.12.2.1.2 1.12.2.2 1.12.3 1.12.3.1 1.12.3.1.1 1.12.3.1.2 1.12.3.1.3 1.12.3.1.4 1.12.3.1.5 1.12.3.1.6 1.12.3.1.7 1.12.3.2 1.12.3.2.1 1.12.3.2.2 1.12.3.3 1.12.3.3.1 1.12.4 1.12.4.1 1.12.4.2 1.12.4.3 1.12.4.4 1.12.4.5 1.12.4.6 1.12.4.7 1.12.4.8 1.12.4.9 1.12.5 References

1.12.1

Ligands featuring covalently tethered moderate to weakly coordinating anions Systems featuring proximal sulfonate anions Phosphine–sulfonate ligands Palladium(II) complexes Nickel(II) complexes NHC–sulfonate ligands Borates and aluminates Scorpionates and related bidentate systems Polypyrazolyl borates Other N-donor borates P-Donor scorpionates and related bidentate systems N- and S-Donor scorpionates and related bidentate systems Scorpionates and related bidentate systems featuring NHCs Mixed NHC borates Mixed Cp systems Diimines and related systems b-Diketiminate platforms and related systems Miscellaneous ligand platforms Monodentate ligands with proximal and distal borate ligand substituents Mono-NHC borates closo-Boron and carborane anions as ligand substituents closo-Dodecaborates as ligand substituents Ligands featuring proximal B12 cages Ligands with distal closo-dodecaborate substituents: Porphyrins, phthalocyanines, and cage complexes Multifunctionalized closo-dodecaborate in ligands closo-Carborane anions as ligand substituents (Alkynyl/acetylide ligands) closo-Carborane anions as ligand substituents for phosphines Phosphine ligands in catalysis N-Heterocyclic carbene ligands with anionic carborane substituents Concluding remarks

373 375 375 375 377 377 378 378 378 382 384 393 393 396 399 399 400 401 402 403 407 407 407 407 408 409 409 409 411 414 416 416

Ligands featuring covalently tethered moderate to weakly coordinating anions

Many metal-mediated catalytic reactions involve the generation of organometallic species that feature an accessible coordination site. Fundamental reactions such as migratory insertion, olefin metathesis, cross-coupling, and hydroaddition reactions require the metal to be able to bind a substrate before the key bond making or breaking steps. When the catalyst is cationic, this binding may be hindered by competition with the solvent or counterion, and this problem is typically minimized with weakly coordinating solvents and weakly coordinating anions (WCAs).1 The latter are the launching point for this chapter. WCAs1 typically share one or more of the following characteristics: (1) they contain peripheral electron-withdrawing groups or atoms; (2) they have a delocalized negative charge; (3) they are relatively large ions. Electron-withdrawing groups or atoms serve to inductively stabilize and also screen the negative charge from positively charged cations. Delocalization thermodynamically stabilizes the charge and shifts the charge density to multiple atoms, rendering each site of charge less potent with respect to its electrostatic potential. The size of the anion is also important as more extended delocalization also tempers the nucleophilicity and basicity of the anion. Furthermore, even if the anionic charge is not delocalized in a p-system, larger anions coated with noncharged groups can effectively shield cations from the electrostatic attraction of the charged site (e.g., borates/aluminates and fluorinated phosphates/antimonates). WCAs have allowed for the exploration of the structure and reactivity of low coordinate main group as well as transition metal and f-element species.1–4 The term noncoordinating anion is often used, but is an exaggeration because there is inevitably some interaction with positively charged ions in solvents of low polarity. These weak interactions can occur via binding of the cation to lone pairs, p-systems,

Comprehensive Organometallic Chemistry IV

https://doi.org/10.1016/B978-0-12-820206-7.00159-1

373

374

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

polarized E–H bonds, and can have varying degrees of covalent and ionic character. Contact ion pairs are more likely to occur in more weakly coordinating solvents such as alkanes, arenes, and fluorinated/chlorinated derivatives of these hydrocarbons. In more coordinating solvents (e.g., water, alcohols, ethers, nitriles) it is more likely to observe solvent-separated ion pairs with only weak electrostatic interactions between the solvated ion pairs and neighboring ions. Some of the most common traditional weakly coordinating anions (e.g., BR4 (R ¼ alkyl or aryl), BF4−, PF6−, SbF6−, Sb2F11−) (Fig. 1A) are formed from the interaction of strong Lewis acids (e.g., BR3, BF3, PF5, SbF5) with alkyl, aryl, or fluoride anions. Alternatively, traditionally weakly coordinating anions are formed from highly oxidized species (e.g., ClO4−, SO2− 4 ) whose coordinative abilities can be further tempered by the introduction of electron-withdrawing groups (e.g., (MeC6H4)SO3−, CF3SO3−, FSO3−) (Fig. 1B). These first-generation WCAs served the pioneers of organic and organometallic chemistry well, exemplified by Olah’s work5 on observing and isolating carbocations in superacid media and the early cationic Rh hydrogenation catalysts developed by Schrock and Osborn.6 However, these anions have limitations with respect to tunability and resistance to decomposition. For example, early metal metallocene species and other low coordinate cations can abstract fluorides from these WCAs or bind very strongly to the oxygen atoms of the oxoanions.3 Many of these anions are also susceptible to chemical and electrochemical oxidation/reduction. Other weakly coordinating anions are those based on fluorinated alkoxy aluminates (Krossing-type anions)7 and fluorinated tetraarylborates such as [BArF4]− (B[3,5-(CF3)2C6H3]4− ¼ Kobayashi’s anion)8 and tetrakis(pentafluorophenyl)borate B(C6F5)4− (Stone’s Anion, BF20)9 (Fig. 1C). These massive anions have their charge effectively screened by fluorinated organic groups. They also have the advantage of rendering their salts soluble in solvents like dichloromethane, arenes, and sometimes alkanes, all of which are weakly coordinating solvents that allow for a clearer picture of the cation’s true reactivity. While the best WCAs are compatible with reactive cations, they may be degraded by carbocations or strong acids via oxidation or bond scission via electrophilic abstraction.3 In addition, the modularity of the syntheses is limited. Another modern class of weakly coordinating anions are the 10- and 12-vertex closo-borane and carborane anions (Fig. 1D).4,10 These species are highly tunable, as the BdH vertices can selectively be substituted with many substituents. When decorated with halogens, these are among the most stable WCAs that are known, as exemplified by the elegant work of Reed and coworkers.3 They are also more resistant to oxidative/reductive decomposition than their 4-coordinate borate cousins.11 These are based on the closo-dodecaborate anion [B12H12]2−, synthesized by Hawthorne in 1960.12 Shortly after, Knoth13 discovered both the 10- and 12-vertex closo-carborane anions. These species are three-dimensionally aromatic, imparting a delocalized electronic structure.14

Fig. 1 (A) Examples of common moderate to weakly coordinating borate anions. (B) Examples of common moderate to weakly coordinating oxoanions. (C) Kobayashi’s and Stone’s Anions. (D) Parent closo-deltahedral hydrido-carborane and borane anions.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

375

These polyhedral anions can be substituted at the BdH and CdH vertices.14 Functionalization via halogenation is readily achieved through electrophilic substitution with elemental halogen reagents or sources. Alternately, nonhalogen functionalization is accomplished through hydroxylation with hydrogen peroxide to form the completely hydroxylated clusters.15 These hydroxyl groups can undergo reactions analogous to alcohols, forming ethers, esters, and carbamate linkages.15 With carborane anions,14 the sp-hybridized CdH vertex can be deprotonated with strong bases and subsequently functionalized with electrophiles, giving them greater synthetic modularity than their all-boron siblings. The use of weakly coordinating anions in organometallic chemistry is a vast topic. This chapter focuses on the applications of moderate to weakly coordinating anions that are covalently tethered to ligand frameworks, and describes their applications in synthesis, small molecule activation, and catalysis. Often these tethered anions are hemilabile, meaning they can dissociate/ reassociate in the presence of a more coordinating substrate or solvent molecule, which renders the ensuing complexes masked low-coordinate species. In one sense, tethered WCAs are no longer “weakly coordinating.” However, the anionic centers in these molecules are held away from the metal center at a specific distance, which can affect their stability or rate of reactions. A survey of ligands and their complexes featuring 4-coordinate borate substituents is covered, followed by systems with aluminates, sulfonates, and closo-deltahedral borane and carborane anions. Rather than aiming to be comprehensive, this chapter highlights the most interesting and important organometallic and related systems possessing such ligands.

1.12.2

Systems featuring proximal sulfonate anions

Most sulfonates are best described as “moderately” coordinating, meaning that they are not easily displaced by most common solvents and mildly nucleophilic substrates. This distinguishes them from “weakly” coordinating groups that are easily displaced by most common solvents, weakly nucleophilic substrates of interest, and often in dynamic equilibrium in solution. Of course there is no clear dividing line in the continuum of coordinative ability. In this section, we do not focus on the plethora of ligands featuring sulfonate groups distant from the coordination sphere of the metal, and instead focus solely on systems featuring proximal sulfonate substituents that are close enough to the metal center that they can behave as hemilabile X-type ligands.

1.12.2.1

Phosphine–sulfonate ligands

The most common type of ligand featuring a proximal sulfonate group are the phosphine sulfonates. Functionalizing aryl groups with phosphine and sulfonate in ortho positions makes a bidentate ligand that may form chelating metal complexes.

1.12.2.1.1

Palladium(II) complexes

Palladium complexes of phosphine–sulfonate ligands have been investigated as catalysts for olefin polymerization reactions. The first successful effort was from Drent and coworkers, where a Pd catalyst was formed in situ by combining Pd(OAc)2 with 1 (Fig. 2).16 This catalyst could copolymerize ethylene with alkylacrylates for the first time. The same catalyst was also used to copolymerize ethylene and CO to obtain a high molecular weight nonalternating copolymer. These complexes were first fully characterized by Reiger and coworkers.17 In addition, they explored the catalytic activity of these complexes for the copolymerization of CO and ethylene, showing that the ensuing polymers contained 30% more CO/ethylene insertion sites compared to Drent’s initial report. Mecking et al. reported the synthesis of water-soluble Pd(II) phosphine–sulfonate complexes and investigated their catalytic activity toward the polymerization of ethylene.18 Later, Nozaki and coworkers reported an anionic methyl Pd(II) catalyst 2 with a phosphine–sulfonate ligand (Fig. 3).19 The copolymerization of ethylene and methyl acrylates was achieved with up to 16% acrylate incorporation. When activated with the

Fig. 2 Drent’s phosphine-sulfonate ligand.

Fig. 3 Nozaki’s anionic Pd(II) catalyst.

376

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

chloride scavengers NaB(ArF)4 or AgOTf, the activity of the system was increased significantly at the expense of selectivity for polar monomer incorporation. This is likely because generating a more accessible coordination site leads to more frequent chain transfer via b-hydride elimination. Furthermore, they reported the use of Pd(II) phosphine–sulfonate complexes for ethylene and acrylonitrile copolymerization.20 The authors could detect the nitrile units in the linear polyethylene backbone using 13C NMR spectroscopy. The same group reported the first example of alternating copolymerization of CO and vinyl acetate using these Pd catalysts, which is notable as a nonradical pathway for synthesizing such polymers (Fig. 4).21 Additionally, they used similar catalysts for the coordination–insertion polymerization of allyl monomers.22 The authors proposed that the chelating sulfonate inhibits b-hydride elimination by favoring intermediates that do not have a coordination site cis to the propagating alkyl chain, thus retarding chain transfer and branch formation. Jordan and coworkers reported a neutral catalyst variant 3 by replacing the chloride with a pyridine ligand (Fig. 5).23 This catalyst was employed in the copolymerization of ethylene and alkyl vinyl ethers. Subsequently, they used similar catalysts for ethylene polymerization and polyethylene incorporated with low levels of vinyl fluoride.24 Later, Jordan and coworkers developed phosphine–sulfonate-based Pd(II) catalysts that produce polyethylene with a higher degree of vinyl fluoride incorporation.25 A significant development in this field was elucidating the influence of the additive B(C6F5)3, which dramatically increases the rate of olefin insertion/chain transfer and thus increases the rate of ethylene oligomerization.26 This was explained by a weakening of the PddO bond when B(C6F5)3 coordinates to one of the oxygens on the sulfonate group, and highlights the differences in reactivity one can achieve by modulating the coordinative ability of a covalently linked anion. The Lewis acid B(C6F5)3 also serves as a scavenger of the L-type ligand, which creates a vacant coordination site adjacent to the methyl group and propagating polymer chain, which is necessary for effective catalysis. The mechanism of copolymerization of ethylene and acrylates by these phosphine sulfonate catalysts was studied in detail, and Nozaki identified two key intermediates that are in equilibrium: the cis-s-complex and trans-p-complex (Fig. 6). The cis-s-complex is lower in energy than the trans-p-complex but the rate determining step (olefin insertion) occurs from the trans-p-complex. Nozaki investigated the energetic differences between analogous cationic (PdP ligated) and neutral (PdO ligated) cis-s-complexes and trans-p-complexes. The neutral cis-s-complex is destabilized relative to cation, which renders the heteroatom in the incorporated monomer chelate more labile. In addition, the trans-p-complex is stabilized relative to its cationic congener, which Nozaki hypothesizes is due to the weak trans-influence of the sulfonate group, resulting in increased p-backbonding to the olefin. The change in charge of the complex from cationic to neutral also likely enhances the backbonding to the bound olefin.

Fig. 4 Nozaki’s alternating copolymerization of CO and vinyl acetate.

Fig. 5 Jordan’s neutral Pd(II) catalyst.

Fig. 6 The phosphine-sulfonate cis-s and trans-p complexes.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

377

To enhance the rate of polymerization, Mecking and coworkers replaced the pyridine in the Jordan system with a more labile ligand like DMSO, which encourages the metal to bind olefins.27 He investigated the use of phosphine sulfonate Pd(II) complexes with a DMSO ligand for the copolymerization of ethylene with different industrially important vinyl monomers, such as vinyl sulfones, acrylates, acrylamides, and acrylic acid.28–34 Subsequently, they showed that this system is also effective for the copolymerization of ethylene with vinyl chloride to produce polyethylenes with a chloride backbone.35 Later, Rieger and coworkers reported the use of the same Pd(II) system for the copolymerization of ethylene with trifluoropropene to produce linear polymers with trifluoromethyl substituents.36 As demonstrated in the section above, the incorporation of a hemilable sulfonate ligand substituent dramatically alters the behavior of Pd(II) olefin polymerization catalysts. These systems allow for the synthesis of novel polymer architectures incorporating polar monomers, which are difficult to achieve with other ancillary ligands or traditional early metal catalysts. The sulfonate ligands retard b-hydride elimination and also control branching via chain transfer and walking. While the reasons behind this behavior are not understood, it is clear that the nature of the pendant anion and the ability to tune it, in this case by adding Lewis acids to weaken the M–O interaction, are important. These observations have led to further explorations of such systems utilizing other transition metals and ligand scaffolds discussed below.

1.12.2.1.2

Nickel(II) complexes

Nickel complexes with phosphine–sulfonate ligands were also explored. Reiger et al. reported the first phosphine–sulfonate nickel complex that polymerizes ethylene.37 Subsequently, Jordan and coworkers reported analogous Ni-based systems that exhibited comparable reactivity.38 Nozaki et al. then reported an allyl nickel complex with phosphine–sulfonate ligands that produced lowmolecular-weight polyethylene.39 Later, Scott and coworkers developed the polymerization of ethylene to produce linear, highmolecular-weight polyethylene, using a single component nickel catalyst 4 with a phosphine–sulfonate ligand (Fig. 7).40 Chen and coworkers reported the use of phosphine–sulfonate nickel complexes for the homopolymerization of norbornene.41 The same group later reported naphthalene-bridged phosphine–sulfonate nickel complexes that showed higher activity for the polymerization of ethylene.42 Furthermore, incorporating a polyethylene glycol unit onto some Pd and Ni catalysts enhanced the rate of ethylene polymerization.43 They reported strategies to synthesize phosphine–sulfonate nickel complexes with high thermal stability and enhanced performance toward the copolymerization of ethylene with polar monomers.44 More recently, they reported Pd and Ni phosphine–sulfonate complexes functionalized to ferrocene that display redox switchable polymerization activity.45 The oxidized catalyst showed decreased performance for ethylene polymerization but was highly active for norbornene polymerization. Stepnicka and coworkers reported similar phosphine–sulfonate ligands functionalized to ferrocene.46 The Rh(I) complexes of these ligands performed catalytic hydroformylation of olefins. Although there are not enough reports to justify a separate subsection, it should be noted that Ziegler and Piers have reported computational and experimental work dealing with the olefin polymerization behavior of b-diketiminate ligands with sulfonate substituents on Pd and Ni.47–49 These studies demonstrate the importance of charge on ethylene copolymerization with polar comonomers, more specifically that reduced overall charge on complexes tends to promote the binding of the polar monomer via the double bond and not the lone pairs of electrons, which is necessary for productive chain propagation. Additionally, a recent report by Carrow demonstrates a thioether–sulfonate chelate Pd(II) system for the catalytic dimerization of thiophenes.50

1.12.2.2

NHC–sulfonate ligands

Metal complexes incorporating N-heterocyclic carbenes (NHCs) have been widely explored in the past few decades due to their ease of preparation and increased thermal stability.51 Like the Pd complexes with chelated phosphine–sulfonate ligands, NHC complexes with tethered sulfonates have also been reported. Nozaki et al. reported the synthesis of the first NHC-sulfonate Pd(II) complexes 5 (Fig. 8).52 Wang et al. reported a series of Pd NHC–sulfonate complexes that showed high catalytic activity for the polymerization of norbornene. Following this, the same group later reported a series of Ru complexes with chelated NHC–sulfonate ligands 6 and investigated the use of these complexes as catalysts for the Ring Opening Metathesis Polymerization (ROMP) of norbornene (Fig. 9).53 Hoveyda et al. reported a NHC–sulfonate Ag complex that can transfer the NHC to Cu. This complex has a chiral NHC that promotes the catalytic asymmetric conjugate addition of dialkyl zinc and trialkyl aluminum reagents to cyclic g-ketoesters.54 This methodology was successfully applied by Hoveyda and coworkers in the total synthesis of complex natural products.55,56 Subsequently, Jordan and coworkers reported a similar Pd sulfonate-NHC complex.57

Fig. 7 Scott’s Ni(II) catalyst.

378

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

Fig. 8 Nozaki’s NHC-sulfonate Pd(II) catalyst.

Fig. 9 ROMP with an NHC sulfonate Ru complex.

1.12.3

Borates and aluminates

The largest group of ligands that feature moderate to weakly coordinating anions in their design are those that feature a borate anion that bears charge, but does not coordinate to the metal since it has no lone pair. The analogous aluminates are less studied, and will be described as well.

1.12.3.1

Scorpionates and related bidentate systems

The study of borates in ligands emerged in 1966 when Trofimenko et al. reported a series of molecules comprised of tetracoordinate boron centers appended with varying numbers of pyrazolyl subunits.58 The reported series of compounds, originally coined “Boron-Pyrazole Chemistry,” was the beginning of a vast number of tetracoordinate borates as ligands in organometallic chemistry (Fig. 10A).59–61 Today, Tp and related systems are common ancillary ligands in organometallic chemistry whose scope has widened beyond N-donors to feature a variety of P, C, O, and S chelates.62–66 The following section will cover the evolution of tetracoordinate borates from their conception to present day with an emphasis on their properties as ligands, as well as notes on their applications to catalysis.

1.12.3.1.1

Polypyrazolyl borates

Borates may contain a varying number of pyrazole subunits attached to the borate, and two ligands of this type are especially prevalent in the organometallic literature: tris(pyrazolyl)borate (Tp) and bis(pyrazolyl)borate (Bp) ligands. Tp ligands are comprised of a boron atom decorated with three pyrazolyl (pz) subunits as well as a single R group (Fig. 10B). Tp can be protonated at N without protonolysis of the BdC bond.58 Tp and Bp both tend to form metal chelates through two cis-coordinating pyrazole N-donors followed by a third interaction, either an agostic interaction or ligation via a third pyrazole subunit, which binds the ligand fac to the metal center.60,61 These were the original members of the class of “scorpionates,” because two pyrazole subunits “pinch” the metal center while a third interaction can “sting” the metal. Tp ligands, along with other borates described herein, are described as “homo-scorpionates” as all three ligating subunits consist of the same N-donor resulting in a k3-N,N,N coordination

Fig. 10 (A) General formula, structure, and geometry of tetracoordinate borate ligands (B) tris(pyrazolyl)borate and (C) bis(pyrazolyl)borate anion reported by Trofimenko et al. in 1966.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

379

mode.59 Bp ligands are accordingly classified as “heteroscorpionates” as an R moiety is substituted in place of the third pyrazole subunit (Fig. 10C).61 Many Tp and Bp ligands are functionalized to achieve varying steric and electronic profiles and have been explored as alternatives to their unfunctionalized parent compounds.67–69 Such functionalization of pyrazolylborates may be as simple as the methylation of the C3 position of the pyrazole subunit, or as extensive as perfunctionalization of the 3, 4, and 5 positions of the pyrazole ring. 3,5-Dimethylpyrazolyl derivatives are particularly easy to prepare, and the steric protection of the substituents mitigates the hydrolysis of the BdN bond (Fig. 11). Direct halogenation of pyrazole is a known means of achieving 3, 4, 5-trihalogenated subunits with enhanced electron-withdrawing character and the introduction of bulky substituents on the pyrazole ring has been shown to direct coordination chemistry and reactivity.70,71 Many of these compounds, particularly complexes of Tp∗(tris(3,5-dimethyl-1-pyrazolyl)borate), have proven to be useful synthons in organometallic chemistry toward catalytically active coordination complexes.72–74 Poly(pyrazolyl)methanes and poly(pyrazolyl)amines are neutral analogs to Tp and Bp ligands. 1.12.3.1.1.1 Ligand synthesis Poly(pyrazolyl)borates are readily prepared by heating an alkali metal borohydride with pyrazole. Di-7, tri-8, and tetra-9substituted borates may be accessed selectively via this method at reaction temperatures of 80  C, 180  C, or > 210  C, respectively (Fig. 12, top).58 A more contemporary approach toward the synthesis of pyrazolyl borates involves the lithiation of functionalized pyrazolyl or other nitrogen donor moieties followed by reaction with a boron halide reagent.58 These strategies have enabled the synthesis of borate ligands containing different ligating moieties as well as tetracoordinate borate ligands, which chelate via other heteroatoms such as carbon, phosphorous, oxygen, and sulfur. As monoanionic ligands, they often can be introduced by adding to metal halides, MX2 or MXL, for homo-10 and heteroleptic-11 transition metal complexes (Fig. 12, bottom). 1.12.3.1.1.2 Electronics Based on the ionic model of electron counting, Tp borates are L2X-type six-electron donor ligands and are isolobal to cyclopentadienyl (Cp) (Fig. 13). Poly(pyrazolyl)borates are electronically distinct from many other tetracoordinate borate ligands, described

Fig. 11 1,3-Borotropic shift yields a mixture of isomers in unsymmetrical pyrazolyl borates.

Fig. 12 (top) Thermally selective synthesis of bis- 7, tris- 8, tetrakis- 9, pyrazolyl borates. (i) excess pyrazoline, 80  C (ii) excess pyrazoline, 180  C (iii) excess pyrazoline, >220  C. (bottom) Synthesis of hetero- 10 and homoleptic 11 Tp metal complexes (iv) 1 eq. 8, 1 eq. MX2 reagent or (v) 2 eq. 8, 1 eq. MXL reagent.

Fig. 13 Generic representation of Bp, Tp, and Cp, complexes, respectively.

380

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

Fig. 14 Modular nature of N-donor sub-units as evidenced by a few reported derivatives.

as Ln-type ligands (vide infra), due to a conjugated p system present in each pyrazole subunit. The s-donating ability of the ligand may be tuned via installation of electron donating and withdrawing groups across the pyrazole subunits. These functionalized pyrazolylborate ligands, denoted Tpx and Bpx, are utilized as anionic ancillary ligands to impart enhanced Lewis acidity at the metal center. Substitution at the 3 and 5 positions with strongly electron-withdrawing groups (dF, dCF3, dNO2) typically enhances this property (Fig. 14).67,75 In addition, these modifications increase the steric bulk of the ligand and when fluorinated substituents are utilized, the ligand can be more resistant to decomposition via intramolecular reactions with the metal center (CdH insertions of carbenoids, cyclometalation, etc.). The ensuing zwitterionic complexes are more electron rich than isoelectronic cyclopentadienyl complexes, as evidenced from the observed v(CO) stretching frequencies of CO complexes.65,76 In contrast, complexes utilizing untethered weakly coordinating anions as a means of stabilizing a formally cationic metal center demonstrate greater Lewis acidic character than those produced using covalently linked borate moieties.

1.12.3.1.1.3 Coordination chemistry/reactivity The robust chemical nature of poly(pyrazolyl)borates is demonstrated by their ability to be purified by sublimation, and stability under acidic, basic, aqueous, and aerobic conditions.58 Tp ligands without great steric bulk are prone to form chemically stable, monomeric, homoleptic complexes (Tp2M) with many transition metals, and the stability of these products is reminiscent of the corresponding metallocenes (Cp2M). Complexes of Bp ligands are distinguished by their tendency to distort into a boat-like conformation, positioning H nuclei in proximity to the chelated metal’s vacant coordination site, and can exhibit agostic interactions for many pendant hydrocarbon functionalities (Fig. 15). Both Tp and Bp ligands have demonstrated significant utility as ancillary ligands to synthesize isolable reactive intermediates and complexes capable of a number of catalytic transformations.62 To date, compounds employing Tp and Bp ligands have been reported for all transition metals generally as an alternative to formally cationic metal centers that employ weakly coordinating counteranions and half-sandwich CpML complexes. While of fundamental interest, homoleptic complexes of poly(pyrazolyl) borates rarely exhibit any catalytic behavior. In contrast, heteroleptic complexes of Tp and Bp borates commonly form chelates with many metal reagents and are employed in a number of catalytic transformations.61 Group 8 and 9 metal complexes utilizing Tp ligands commonly display alkene and alkyne oligomerization chemistry and mediate a number of coupling reactions as CdH activation catalysts.77,78 Particularly notable are a series of manuscripts from Jones and coworkers that utilize RhTp complexes, which are analogous to well-studied Cp Rh fragments,79 to probe alkane CdH oxidative addition/reductive elimination. One such investigation describes evidence for an associatively induced reductive elimination of benzene from a RhTp complex illustrating one-way Tp and Tpx ligands have been utilized to investigate fundamental reaction mechanisms in organometallic chemistry (Fig. 16).80–89 This example also demonstrates the importance of substitution of the Tp ligand to increase steric bulk in the coordination sphere of the metal, which in this case aids in the partial dissociation of a scorpionate arm to accommodate the newly formed coordinated benzene molecule. Among the vast number of these Tpx ligands, a handful have been given particular attention for their unusual coordination chemistry and ability to mediate catalytic transformations, and provide mechanistic insight into enzyme behavior.90,91 3,5-Diisopropyl-2-pyrazolylborate (TpiPr2) is one such ligand which is reported to form heteroleptic complexes upon reaction with MX2 (M ¼ Zn, Mn, Fe, Cu, Co).61 The tetrahedral complex TpiPr2MCl has been used as an synthon toward a number of dinuclear (OH)2, CO3, and O2-bridged metal complexes.68,92,93 The side-on bridged dioxygen complex (TpiPr2Cu)2O2 was the first spectroscopic and structural mimic for dioxygen binding in hemocyanin, which benefited from the similarity of the pyrazole donors to the natural histidine imidazoles.94 The kinetic stability of this molecule can be explained by the fac coordination of the ligand, which projects the bulky iPr groups around the metal center like a crown, kinetically protecting the coordinated O2 from further reduction by a third metal fragment. Similarly, the smaller TpMe borates have been employed for structural studies of Cu complexes to offer mechanistic insight into a variety of enzymes.95

Fig. 15 Generic representation of pseudo-boat distorted Bp complex displaying an agostic interaction.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

381

Fig. 16 The reductive elimination of an aryl group from a RhTp complex to generate benzene.

Fig. 17 Cationic group 4 complexes of bis-pyrazolyl borate yield highly active ethylene polymerization catalysts.

Toward catalytically active coordination complexes, Tp and Bp borates functionalized with sterically bulky alkyl or aryl groups such as TpMes or BpMes have led to a number of group IV complexes which are very active ethylene polymerization catalysts (Fig. 17).96 High catalytic activity of these complexes is attributed to the stabilization, relative to complexes with untethered WCAs, of cationic Zr or Hf metal centers by the formally anionic borate ligand.74 Some Tpx ligands have been functionalized with strong electron-withdrawing functional groups.97,98 3,5-bis-trifluoromethyl2-tris(pyrazolyl)borate (TpCF3) is one such ligand that has demonstrated greatly enhanced electrophilicity of metal centers in tetrahedral TpCF3AgL complexes.99,100 In contrast to their unfunctionalized Tp counterparts, which often decompose via metal carbene and nitrene insertions into the CdH bonds, these “Teflon-coated” TpCF3AgL complexes have been reported to mediate carbene insertion reactions into CdX and aryl CdH bonds (Fig. 18), as well as the most challenging hydrocarbons like methane.101 In general, other Tpx borates functionalized with strong electron-withdrawing moieties have been shown to catalyze similar molecular transformations. Cu complexes of 3,4,5-tribromo-2-tris(pyrazolyl)borate (TpBr3CuL), TpCF3CuL, and BpCF3CuL have all been reported to catalyze nitrene insertion into arene CdH bonds to achieve the amination of a small library of sp2 and sp3 hydrocarbons (Fig. 19).69,102,103 Interestingly, the BpCF3CuL complex shows higher yields in arene amination products. It is postulated that reduced steric bulk as a result of one fewer 3,5-bis-trifluoromethylated pyrazole ring, as well as a low energy barrier associated with freeing an additional coordination site on the metal, promotes nitrene addition and contributes to an overall greater yield (Fig. 19C). For this reason, Bp and Bpx borate ligands are especially intriguing given their potential for k3 N,N,H coordination via an agostic interaction from the pendant R group. Cu complexes of Bp borates have also been investigated for their ability to mediate carbene insertion into NdH bonds with some success,104,105 but they are not as effective in catalysis compared to their fluorinated tris(pyrazolyl)borate cousins (Fig. 19B). These applications and more, as well as their coordination chemistry, structure, and bonding, and Tp and Bp metal chelates have been reviewed multiple times over the past few decades as tetracoordinate borate ligands continue to demonstrate their value toward pursuits in organometallic chemistry.59,61,62,106 1.12.3.1.1.4 Related bidentate and tridentate aluminates Weakly coordinating anions based on aluminum have recently seen increased interest. This is particularly the result of groundbreaking work by Krossing and coworkers.2,7,107,108 Storr109 and coworkers reported the first bona fide example of an aluminate

382

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

Fig. 18 Examples of carbene insertion reactions utilizing a TpCF3Ag(THF) complex.

Fig. 19 (A) Perbrominated tris(pyrazolyl)borate Cu complex (B, C) 3,5 substituted trifluoromethylated bis- and tris(pyrazolyl)borate complexes.

complex analogous to pyrazolyl borates. In this report the authors reacted triethylaluminum with a mixture of sodium pyrazolide and pyrazole to generate [Al(pz)4]Na. The authors then attempted to create simple chelated transition metal complexes to no avail. The first X-ray structure of a poly(pyrazolyl)aluminate complex was reported in 2005.110 In 2011, Winter111 and coworkers isolated a variety of aluminate hydride complex with the general formula [Al(R2pz)nH(4-n)][Li(THF)x]. Perhaps the paucity of reports of such complexes is due to the more labile AldN bond compared to the BdN bonds in Tp and related ligands.

1.12.3.1.2

Other N-donor borates

A significant disadvantage of pyrazolyl-based N-donors is the tendency for the pyrazolyl borate BdN bond to undergo hydrolytic cleavage and 1,3-borotropic isomerization (Fig. 11) of the ligand. Popular strategies to mitigate this issue involve functionalization of pz subunits at the 3 and 5 positions to achieve a symmetrical subunit. A variety of tetracoordinate borates that chelate metals via N-donor atoms, and do not introduce sterically demanding Tpx or Bpx pyrazole derivatives, exist which instead utilize pyridine, imidazole, triazole, and oxazole moieties in place of pyrazole (Fig. 20A–D).

Fig. 20

(A) Tris-2-(pyridyl)borate ligand, (B) tris(imidazolyl)borate ligand (c), tris(1,2,4 triazolyl)borate ligand, and (D) tris(4,4 dimethyl-2-oxazolinyl)borate ligand.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

383

Fig. 21 The oxidation of cyclohexane mediated by Tox complex 12.

Tris(oxazoline)borates (Tox) are one such alternative to pyrazolyl ligands that mitigate some of these shortcomings.112 With this ligand, oxazole subunits are covalently tethered to the boron nuclei via BdC linkages. The lack of steric congestion around the metal centers of Tox coordination complexes does not hinder catalytic activity as has been postulated with sterically bulky Tpx ligands. To this end, heteroleptic (Tox)CoX complexes, such as 12, have been reported to act as catalysts for the conversion of C6H12 to cyclohexanol without excessive quantities of overoxidation products (Fig. 21).113 Cobalt complexes of To have been exploited for alkane functionalization and catalytic oxidative carbonylation chemistry. Along with To ligands, diphenyl-bis(oxazolyl)borate ligands (Borabox ligands, Box) have received a large amount of attention as anionic analogs to the widely reported carbon-centered bis-oxazolines (Box). Box ligands are known to stabilize metal centers capable of mediating asymmetrically catalyzed reactions through a combination of electronic and steric effects.114 In addition, Box ligands, such as (ArF)2B(BotBu) 13 (Fig. 22), have been reported to demonstrate excellent stereoselectivity in cyclopropanation reactions. For example, a copper complex of 13 gives 99% stereoselectivity for the trans isomer in multiple cyclopropanation reactions with 2,6-ditert-butyl-4-methylphenyl diazoacetate115 (Fig. 22). Tris(2-pyridyl)borate (Tpyb) is another N-donor borate ligand that displays significantly different properties from pyrazolyl borates. Pyridyl borates have stronger s-donating ability than pyrazolyl borates and, as with other N-donor alternatives, are also appended to the borate nuclei via a BdC covalency.116 As “third-generation” borate ligands of Tp and Bp, their applications as ligands have only recently begun to be explored. Ru complexes of Tpyb borates have been reported to be highly active ammonia fixation catalysts and a convenient means of producing borazine from boramine.117,118 Tpyb ligands are among the series of tetracoordinate borates that are being explored as ligands in coordination polymers for materials science applications.119 Additionally, tetrazolylborate ligands present themselves as an attractive substitute to traditional Tp ligands. Tetrazolylborates exhibit increased electron-withdrawing character due to its additional nitrogen atoms. In some circumstances, tetrazolylborates can form coordination polymers, ligating the metal center at both N2 and N4. These compounds demonstrate appreciable water solubility (Fig. 23).71

Fig. 22 Ligand 13 and its ability to induce high stereoselectivity in Cu catalyzed cyclopropanation chemistry.

Fig. 23 2D Coordination polymer based on the nitrogen-rich (bis)tetrazolyl-borate ligand.

384

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

The BdN linkage in their design renders them susceptible to hydrolysis and isomerization of the ligand. Like their Tp analogs, octahedrally coordinated bis-chelates result from complexation of tetrazolylborates to transition metals. Additionally, the increased number of N-donor atoms present on these ligands tend to result in complex coordination environments and extended 2D and 3D metal chelates.120 Tetrazolyl borates have not commonly been investigated for their catalytic behavior and receive more attention in the field of materials science toward the synthesis of coordination polymers. 1.12.3.1.2.1 Heteroscorpionates Tetracoordinate borate ligands need not consist solely of two or three of the same N-donor subunit. Within the past decade, efforts directed toward exploring heteroscorpionates, or scorpionate ligands that consist of multiple different donor subunits covalently tethered to a single boron, have been explored as ligands. One such example is [HB(pz)2Py]−, which is comprised of two pyrazolyl subunits and a single pyridyl N-donor. Heteroscorpionates demand more involved synthetic methods to produce in contrast to the one-pot thermally driven synthesis of Tp and Bp ligands. As a result, these ligands are the subject of less study regarding the catalytic activity of their metal chelates.

1.12.3.1.3

P-Donor scorpionates and related bidentate systems

Bis and tris(phosphino)borates have been established as reliable anionic phosphine ligands since their introduction at the turn of the millennium (Fig. 24).121,122 They generally undergo faster CdH activation reactions compared to cationic analogs.76 Furthermore, many (phosphino) borates have yielded metal complexes capable of carbene, nitrene, and atom transfer reactions.103 Ligand synthesis, electronic properties, and interesting coordination chemistry and reactivity are discussed below. 1.12.3.1.3.1 Ligand synthesis Bis(phosphino)borate ligands were synthesized in two steps. In step 1, the tetraalkyl tin reagents were treated with BCl3 to make dialkylchloroborane reagents (Fig. 25).64 In the second step, the dialkylmethyl phosphines were treated with BuLi in TMEDA to deprotonate the methyl to give the anionic phosphine. Two equivalents of this anionic phosphine were treated with the dialkylchloroborane from step 1 to produce the desired ligand after losing two equivalents of lithium chloride. Sometimes the deprotonation of the phosphine was more productive when protected with BH3, especially when aryl groups like p-trifluoromethylphenyl were present. The BH3 group could be removed after the ligand synthesis by heating in morpholine at 60 C for aryl phosphines, while BH3-protected alkylphosphino derivatives could not be deprotected. Tris(phosphino)borates are synthesized via the reaction of 3 equivalents of lithiated phosphine and the desired BRCl2 fragment. 1.12.3.1.3.2 Electronics Similar to Tp systems, bi- and tridentate phosphinoborates are considered four-electron LX and 6-electron L2X-type ligands, respectively. Unlike Tp ligands, which feature a conjugated pi-system that induces electronic effects on the donating ability of the nitrogen, phosphinoborates feature a methylene spacer, which decouples electronic effects induced by the borate nuclei from the phosphine. Peters et al. investigated the donating ability of a series of bis-phosphinoborate ligands by observing the resulting CO stretching frequency (nCO) of a Pt complex with variable boron and phosphine substituents (Fig. 26).64 The s-donating ability of phosphinoborates may be tuned by directly substituting the phosphine with electron donating or withdrawing moieties. Consequently, substitution undertaken directly on the borate nuclei induces minimal electronic effect on the metal center and serves

Fig. 24 First reported bis- and tris-phosphino borates.

Fig. 25 Synthesis of bis(phosphino)borate ligands.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

385

Fig. 26 nCO stretching frequencies of a series of bis(phosphino)borates ligands.

primarily as a means of influencing the steric properties of the ligand. Infrared spectroscopy of Rh(CO)2 fragments bearing TpMe2, Cp∗, and PhBPPh3 confirmed the tris(phosphino)borate to be the strongest donor.123 1.12.3.1.3.3 Reactivity and coordination chemistry The first examples of coordination complexes bearing these ligands were introduced by Tilley et al. via reaction of [Li(TMEDA)] [PhB(CH2PPh2)3] (TMEDA ¼ N,N,N,N,-tetramethylethylenediamine) with [(COE)2IrCl]2 to yield (PhBPPh3)IrH(3-C8H13), 14, as confirmed by multinuclear NMR and X-ray crystallography (Fig. 27).121 Initial exploration of the reactivity of 14 reactivity with H2SiMes2 at elevated temperatures in benzene formed silylene complex [PhBPPh3](H)2Ir]SiMes2 15, illustrating the direct conversion of a silane to a silylene fragment at a transition metal center. A later report expanded the scope to include more secondary silanes and the synthesis of an analogous germylene complex [PhBPPh3] (H)2Ir]GeMes2. Furthermore, primary silanes reacted with 14 to produce silylene complex (PhBPPh3)H2Ir]SiHR, which reacted further in the presence of cyclooctene (COE) to produce (PhBPPh3)H2Ir]SiR(c-C8H15) 16 (R ¼ Mes or 2,4,6-triisopropylphenyl) (Fig. 27). Tertiary silanes reacted with 14 to yield silyl-capped trihydride complexes of the type [PhBPPh3]IrH3(SiR3) (R ¼ Et; R ¼ Me).123 (PhBPiPr3)IrH(3-C8H13) 17 could be synthesized in a manner analogous to 14 and led to the discovery of new reactivity with secondary silanes, yielding silyl-capped trihydride complexes 18.123 In contrast, when 17 was exposed to secondary silanes at elevated temperatures the product was silyl-capped trihydride complexes 18 (Fig. 28). Complex 17 reacted with neutral donor ligands to give [PhBPiPr3]IrH2L (L]PMe3, PH2Cy, and CO) and releasing 1,3-cyclooctadiene. If [PhBPiPr3]Li(THF) is reacted with {(COE)IrCl}2 in a propene-saturated THF solution, the allyl analog of 17, [PhBPiPr3]IrH(3-C3H5), was isolated. Unlike complex 17, [PhBPiPr3]IrH(3-C3H5) does not react with Et2SiH2 or Ph2SiH2, due to the slower b-hydride elimination.

Fig. 27 Synthesis and reactivity of novel complex 14 with silanes.

Fig. 28 Reactivity of complex 17 with secondary silanes.

386

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

Fig. 29 Synthesis of complexes 19 and 22.

The allyl complexes 14 and [PhBPPh3]Ir(H)(3-C3H5) 20 were utilized as precursors for dihalides [PhBPPh3]IrI2 19 and [PhBPPh3] IrCl2 21 (Fig. 29).124 Chlorination at the boron nuclei was observed at temperatures above 90 C during the synthesis of 21. 19 was reacted with two equivalents of methyllithium to generate [PhBPPh3]IrMe2. 14 and 20 both react with CO to give [PhBPPh3]Ir(CO)2. Reacting complex 20 with molecular hydrogen in benzene provides {[PhBPPh3]IrH2}2. Compound 20 catalyzes H/D exchange of COE under D2 in benzene-d6. Evidence of H/D exchange at positions ortho to phosphorus was demonstrated with reactions using D2SiMes2. Further evidence is demonstrated via isolation of cyclometalated product 23 via the reaction between 14 and PMe3 (Fig. 30). The complex [k2-PhBPiPr3]Rh(PMe3)2 24 can be produced from reacting [PhBPiPr3]Li(THF) with (PMe3)4RhOTf. Complex 24 was found to undergo a dynamic equilibrium process involving dissociation of the PMe3 ligands and reversible migration of a methylene group in the ligand backbone from B to Rh at elevated temperatures (Fig. 31). The ensuing product is [PhB(CH2PiPr2)2] RhCH2PiPr2 25, and can be independently generated from reacting [PhBPiPr3]Li(THF) with [RhCl(C2H4)2]2. Complex 24 also reacts with CO to generate [PhBPiPr3]Rh(CO)2. Additionally, complex 24 undergoes oxidative addition with molecular hydrogen to generate [PhBPiPr3]RhH2(PMe3). Compound 25 reacts with Ph2SiH2 to generate [PhB(Ch2PiPr2)2]RhH2(SiHPh2)PMe3 26 with loss of a ligand arm (Fig. 31). The complex [PhBPPh3]Rh(COD) 27 was an efficient catalyst for the selective hydrogenation of a,b-unsaturated aldehydes.125 27 was synthesized by metathesis of {Rh(m-Cl)(COD)}2 with [Li(TMEDA)][PhBPPh3] in dichloromethane. 27 released cyclooctane when exposed to hydrogen in acetonitrile to produce [PhBPPh3]Rh(H)2(NCMe) which interestingly reacted with chloroform creating CH2Cl2 and an equimolar mixture of the cis and trans isomers of [(PhBPPh3)RhH(m-Cl)]2, which shed its hydrides to yield [(PhBPPh3)Rh(m-Cl)]2, 28 (Fig. 32). Two years later, [PhBPPh3]Rh(H)2(NCMe) was reported to efficiently catalyze the dimerization of aliphatic and aromatic, enolizable, and nonenolizable aldehydes.126 The rhodium complex [PhBPPh3]Rh(CH2]CH2)(NCMe) also reacted with H2 to afford [PhBPPh3]Rh(H)2(NCMe) which in solution slowly dimerizes to {[PhBPPh3]Rh(H)(m-H)}2. Complex [PhBPPh3]Rh(CH2]CH2)(NCMe) was oxidized with [Cp2Fe]+ in acetonitrile to afford [[PhBPPh3]Rh(NCMe)3]+. [PhBPPh3]

Fig. 30 Cyclometallation of 14 via addition of PMe3.

Fig. 31 Reversible migration of a dCH2 group yielding complexes 24 and 25, and subsequent reaction with Ph2SiH2 to yield 26.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

387

Fig. 32 Tp complex 27 and resulting chloro-bridged dimer resulting from hydrogenation of cyclooctadiene and subsequent reaction with chloroform.

Fig. 33 Synthesis of bis(phosphino)borate Rhodium complex 30.

Rh(CH2]CH2)(NCMe) will react with carboxylic acids, such as benzoic or acetic acid, to form complexes of the type [PhBPPh3] Rh(Z1-C2H5)(k2-O2CR) (R ¼ Ph or Me). These ethyl complexes establish an equilibrium with the corresponding hydride compounds [PhBPPh3]Rh(H)(k2-O2CR) with ethylene loss. Zwitterionic bis(phosphino)borate rhodium(I) complexes can be synthesized by reacting [Ph2BPPh2] with [(nbd)Rh(m-Cl)]2 or [(COD)Rh(m-Cl)]2 to generate [Ph2BPPh2]Rh(nbd)], 30, and [Ph2BPPh2Rh(COD)], respectively (Fig. 33).127 30 can mediate the hydrogenation of styrene to ethylbenzene in d6-acetone. The norbornadiene ligand in 30 can be removed via hydrogenation in acetonitrile to generate [Ph2BPPh2]Rh(CH3CN)2, or substituted with excess to PMe3 to generate [Ph2BPPh2]Rh(PMe3)2. The [Ph2BPPh2]Rh(CH3CN)2 complex was shown to catalyze the hydroacylation of 4-methyl-4-pentenal. The [Ph2BPPh2Rh(COD)] complex can be treated with CO gas to give [Ph2BPPh2]Rh(CO)2, which can then undergo single substitution by PMe3, PPh3, and an N-heterocyclic carbene (NHC) to generate [Ph2BPPh2]Rh(CO)(PMe3), [Ph2BPPh2]Rh(CO)(PPh3), and [Ph2BPPh2]Rh(CO)(NHC), respectively. A series of divalent cobalt(II) complexes [PhBPPh3]CoI, {[PhBPPh3]Co(m-Br)}2, and {[PhBPPh3]Co(m-Cl)}2 were prepared and their electronic properties were reported.128 Tl[PhBPiPh3], 31 readily reacts with CoI2 to generate [PhBPPh3]CoI, 32. Attempting to synthesize 32 directly from [Li(TMEDA)][PhBPPh3] and CoI2 yielded ill-defined mixtures, illustrating the primacy of the Tl reagent. Bromide- and chloride-bridged dimers were synthesized in two steps from [PhBPPh3]CoI, starting with the addition of thallium 2,6-dimethylphenolate Tl(O-2,6-Me2Ph) to generate (PhBPPh3)Co(O-2,6-Me2Ph), which was subsequently reacted with NaCl or KBr to yield {(PhBPPh3)Co(m-Br)}2 or {(PhBPPh3)Co(m-Cl)}2, respectively. Exposure to O2 caused insertion of O atoms into two of the three PdCo bonds to generate two new P]OdCo linkages. A benzene solution of 32 reacted with PMe3 to generate [PhBPPh3] Co(PMe3)I, which was then reduced by sodium amalgam to give the pseudotetrahedral cobalt(I) complex [PhBPPh3]Co(PMe3) 33, a useful synthon for further synthetic modifications (Fig. 34).129 Tl[PhBPiPr3] 34 reacted with CoX2 (X ¼ Cl, I) and generated (PhBPiPr3)CoCl, and [PhBPiPr3]CoI, 35.130 Complex [PhBPiPr3]CoCl reacted with CO to produce divalent [PhBPiPr3]CoCl(CO). Interestingly, 35 is reduced in the presence of CO to generate monovalent [PhBPiPr3]Co(CO)2. [PhBPiPr3]CoCl underwent similar oxidation processes to create P]OdCo linkages when exposed to oxygen. A series of complexes (PhBPR3)CoX (X ¼ OR, SR, OSiR3, SSiR3) were synthesized from synthons 32 or 35 and a thallium salt of respective aryloxide, siloxide, thiolate, or silylthiolate. The structural and magnetic properties of the aforementioned series and cobalt(III) complex [(PhBPPh3)Co(OSiPh3)][BAr4] (Ar ¼ Ph or m-(CF3)2Ph) were reported. Addition of two equivalents of (p-tolyl)azide to 33 in benzene yielded the first terminal imido complex of cobalt, [PhBPPh3] Co ⫶N-p-tolyl. [PhBPPh3]Co ⫶ N-p-tolyl was reacted with CO to release p-tolyl isocyanate and cobalt dicarbonyl complex, [PhBPPh3] Co(CO)2. Complex [PhBPPh3]Co(PMe3) also reacted with 2 equivalents of Ph2CN2 to give the diazoalkane complex, [PhBPPh3] Co(N2CPh2) 36 (Fig. 35). Diazoalkane reactivity was explored utilizing bis(phosphino)borate complex [Ph2BPtBu2]Co(2,4,6-trimethylaniline). Products analogous to 36 were observed when [Ph2BPtBu2]Co(2,4,6-trimethylaniline) was reacted with (Me3Si)2CN2 or MesCN2. However, the reaction with Ph2CN2 produced [Ph2BPtBu2]Co complexes bearing terminal carbene or Z1-azine ligands.

Fig. 34 Cobalt complexes prepared via transmetalation from a Tl complex.

388

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

Fig. 35 Formation of diazoalkane complex 36.

Fig. 36 Synthesis of bis(phosphino)borate copper complexes.

Several copper(I) complexes supported by bis(phosphino)borate ligands have also been reported.131 In search of a reliable entry point into copper complexes, Peters et al. began their investigations utilizing 3 different bis(phosphino)borate 37–39 (Fig. 36). When CuI was reacted with [ASN][(p-tBuPh)2BPPh2] 37, the product was [ASN][[p-tBuPh)2BPPh2]CuI]. The reaction between Li(Ph2BPiPr2) 38 or Li(m-xylyl)2B(PtBu2)2 39 and CuI gave immediate degradation. However, the less redox susceptible CuCl yielded the desired (Ar2BPR2)CuCl complexes with the generation of some minor degradation products. One of these degradation products was characterized as neutral borane analog (m-xylyl)B(PtBu2)2CuCl, which is the result of dearylation of the boron backbone. Synthetically useful Cu precursors, (Ph2BPiPr2)Cu(MeCN) 40 and ((m-xylyl)2BPtBu2)Cu(MeCN) 41, were also developed via the salt metathesis reaction of 38 or 39 with [Cu(MeCN)4][PF6] in THF. The labile acetonitrile ligand was substituted with L-type ligands such as PMe2Ph, S]PMe3, 2,6-lutidine, and CNtBu. These procedures were modified to synthesize [(PhBPPh2)Cu(pbb)] [BF4] (pbb ¼ 2-(20 -pyridyl)benzimidazolylbenzene) via the reaction of [PhBPPh2][ASN] with [Cu(CH3CN)4][BF4] in CH2Cl2.132 The electronic and photophysical properties of this complex were examined. In 2012, the complexes [PhBPPh3]Cu(L)] (L ¼ PPh3 or THF) were synthesized and the catalytic properties of [PhBPPh3] Cu(PPh3)] were reported.133 [PhBPPh3]Cu(THF)] was generated by reacting [PhBPPh3]Li(TMEDA) and CuCl in tetrahydrofuran. Similarly, [PhBPPh3]Cu(PPh3) was synthesized by reacting [PhBPPh3]Li(TMEDA) and CuCl in tetrahydrofuran, followed by the addition of PPh3. [PhBPPh3]Cu(PPh3) was found to catalyze olefin cyclopropanation, cyclopropenation, aziridination, halocarbon addition, and atom transfer radical polymerization, OdH and NdH functionalization of carbene insertion, and furan conversion into dihydropyridines (Fig. 37). In 2008, a series of tris(phosphino)borato silver(I) complexes were synthesized and included [PhBPPh3]AgPEt3, [nBuBPPh3] AgPEt3, [PhBPiPr3]AgPEt3, and [nBuBPiPr3]AgPEt3.134 The silver complexes were synthesized via the salt metathesis reaction between AgCl and the corresponding Li(TMEDA) tris(phosphino)borate. In 2014, the fluorinated ligand [PhBPp-CF3Ph3] was developed and used to synthesize [PhBPp-CF3Ph3]Tl, [PhBPp-CF3Ph3]Cu(PPh3), and [PhBPp-CF3Ph3]Ag(PPh3).103 The thallium reagent was generated by reacting [PhBPp-CF3Ph3(BH3)3]Li(Et2O) with TlNO3.

Fig. 37 Cyclopropan−/cyclopropenation catalyzed by 42.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

389

Complex [PhBPp-CF3Ph3]Cu(PPh3) was synthesized by reacting [PhBPp-CF3Ph3]Tl with CuI and PPh3. Complex [PhBPp-CF3Ph3] Ag(PPh3) was synthesized by reacting [PhBPp-CF3Ph3]Li(TMEDA) with AgCl, followed by addition of PPh3. Both the copper and silver complexes were reported as catalysts in nitrene transfer reactions to saturated and unsaturated substrates. Aziridination of olefins with observed with the copper complex, while the CdH bonds of cyclic ethers were amidated with the silver complex as the catalyst. Treating 31 with RuCl2(PPh3)3 gave 28, which is an efficient catalyst for the selective hydrogenation of a,b-unsaturated aldehydes.125 A later report reacted 28 with 2 equivalents of AgPF6 in an acetonitrile/THF solvent mixture to synthesize [(PhBPPh3)Ru(NCMe)3]PF6, a hydrogenation transfer catalyst precursor.135 The following year, the complex [PhBPPh3]RuCl(PMe3) was synthesized by reaction of 31 with RuCl2(PPh3)3.130 Years later, complex [PhBPPh3]RuCl(PMe3) was reported to react with PhSiH3 in THF to form silyl-capped trihydride complexes [PhBPPh3]Ru(m-H)3Si(Ph)Cl(PMe3).136 In 2011, Z3-silane s-complexes [PhBPPh3]RuH(Z3-H2SiPhMe] and [PhBPPh3]RuH(Z3− H2SiPh2] were generated by reacting 28 with excess PhMeSiH2 or Ph2SiH2, respectively.136 Lewis acidic adducts are present on the Si atom evident by the formation of adducts with THF and 4-dimethylamino-pyridine. Three years later, the same group reported the reaction of 28 and (THF)2Li(SiHMes2) to generate [PhBPPh3]Ru[CH2(2-(Z3-H2SiMes)-3,5-Me2C6H2)].137 Complex [PhBPPh3]Ru[CH2(2-(Z3-H2SiMes)-3,5-Me2C6H2) was reacted with H2 to form [PhBPPh3]Ru[CH2(2-(Z3-H2SiMes2), which was found to eliminate Mes2SiMe2 and form [PhBPPh3] Ru(Z5-C6H7). Furthermore, [PhBPPh3]Ru[CH2(2-(Z3-H2SiMes)-3,5-Me2C6H2) was reported to exist in equilibrium with a 16-electron silylene complex [PhBPPh3]Ru(m-H)(¼SiMes2), as evident from DFT calculations and trapping of the silylene with XylNC (Xyl ¼ 2,6-dimethylphenyl) to form [PhBPPh3]Ru(CNXyl)(m-H)(]SiMes2). Reacting FeX2 (X ¼ Cl, Br, I) with 31 produced 4-coordinate high-spin iron(II) complexes [PhBPPh3]FeX. [PhBPPh3]FeCl, 43, was reacted with Me2Mg to generate [PhBPPh3]FeMe.138 Exposure of [PhBPPh3]FeMe to hydrazine formed [PhBPPh3]Fe(Me)(Z2-N2H4), which can be oxidized by [Fc][PF6] to form {[PhBPPh3]Fe(NH3)(Z2-N2H4)}{(PF6)} or reacted with CO to form [PhBPPh3]Fe(CO) (Z2-N2H4).139 If [PhBPPh3]Fe(CO)(Z2-N2H4) is treated with [Fc][PF6], then the cationic hydrazine species {[PhBPPh3]Fe(CO) (Z2-N2H4)}{PF6} is formed. Complex 43 reacted with sodium azide in acetonitrile to give 44 (Fig. 38).140 Complex 44 was reduced with sodium amalgam to generate 45. 45 reacted with CO to generate [PhBPPh3]Fe(CO)2 and [PhBPPh3]Fe(CO)2(NCO) as the major and minor products, respectively. Compound 45 also reacted with H2 to generate [([PhBPPh3]Fe)2(m-NH)(m-H)] [Na(THF)5] 46 (Fig. 39).141 Complex 46 was oxidized by [NO+][PF6−] to yield neutral complex {([PhBPPh3]Fe)2(m-NH)(m-H). In addition, [PhBPPh3] Fe(CO)2Cl was generated from 43 and CO in benzene.140 Reduction of 43 by sodium/mercury amalgam in THF in the presence of 3 equivalents of PPh3 generated [PhBPPh3]Fe(PPh3).142 The complex [PhBPPh3]Fe(PPh3) can be oxidized by p-tolyl azide in benzene to give [PhBPPh3]Fe^N(p-tolyl) 47 (Fig. 33). Complex 47 reacts with CO to release p-tolyl isocyanate and [PhBPPh3] Fe(CO)2 48 with regeneration of 47 via addition of p-tolyl azide (Fig. 40). Exposure of 47 to H2 generates [PhBPPh3]FeNH(p-tolyl) 49, which slowly reacts further to generate p-toluidine and the partially hydrogenated benzene complex [PhBPPh3]Fe(Z5-cyclohexadienyl) 50 (Fig. 41).143 [PhBPPh3]Fe(PPh3) also reacted with 1-azidoadamantane in benzene to give [PhBPPh3]Fe^N(1-Ad).140 This complex can be reduced by sodium amalgam, followed by cation exchange with [nBu4N][Br] to give [PhBPPh3Fe^N(1-Ad)][nBu4N]. Complex 46 can be reduced by PCl3 to produce the nitride complex {([PhBPPh3]Fe)2(m-N)}.141 [PhBPiPr3]Tl, 34, was prepared by addition of TlPF6 to [Li(TMEDA)] [PhBPiPr3], which reacted with FeCl2 in THF solution to yield [PhBPiPr3]FeCl 51 (Fig. 42). Complex 51 reacted with excess CO to give divalent species [PhBPiPr3]FeCl(CO). The lithium amide reagent Li(dbabh) (dbabh ¼ 2,3:5,6-dibenzo-7-aza bicyclo[2.2.1]hepta-2,5-diene) reacted with 51 to generate [PhBPiPr3]Fe(dbabh) 52 (Fig. 42).144 Upon heating, thermally unstable complex 52 generated the nitride complex [PhBPiPr3]Fe^N 53 following the loss of anthracene (Fig. 42). Complex 53 underwent bimolecular condensation in benzene via nitride coupling to generate {[PhBPiPr3]

Fig. 38 Reactivity of 43 with sodium azide, followed by reduction of 44 with sodium amalgam.

Fig. 39 Reactivity of complex 45 with H2 to yield complex 46.

390

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

Fig. 40 Reactivity of 47 with carbon monoxide to yield 48, which may subsequently be regenerated to 47 with p-tolyl azide.

Fig. 41 Reactivity of 47 with H2.

Fig. 42 The synthesis of the iron nitride complex 53.

Fe}2(m-N2). Complex {[PhBPiPr3]Fe}2(m-N2) reacted with excess PMe3 to generate [PhBPiPr3]FePMe3.145 Complex 51 can react with KEt3BH in the presence of PMe3 to generate [PhBPiPr3]Fe(H)PMe3. MeLi or neopentyllithium can react with complex 51 to give [PhBPiPr3]Fe-Me and [PhBPiPr3]Fe-CH2CMe3, respectively. Complex 51 also reacts with PhCH2MgCl to generate [PhBPiPr3]FeCH2Ph. These alkyl complexes [PhBPiPr3]FedR (R ¼ Me, CH2CMe3, or CH2Ph) can be hydrogenated under 1 atm of H2 with the presence of PR3 (PR3 ¼ PMe3, PEt3, or PMePh2) to generate RH and the iron(IV) trihydrides [PhBPiPr3]Fe(H)3PR3. The reaction of [Li(THF)][PhBPiPr3] with FeBr2(THF)2 forms [PhBPiPr3]FeBr.146 When [PhBPiPr3]FeBr is reacted with KSi(SiMe3)3 in benzene, a silylated benzene moiety is formed in Z5-coordinated [PhBPiPr3]Fe(Z5-6-Si(SiMe3)3C6H6). In contrast, when complex [PhBPiPr3] FeBr is reacted with KSi(SiMe3)3 in pentane, the product isolated is [PhBPiPr3]FeSi(SiMe3)3. Reacting [PhBPiPr3]FeBr with Ph3SiLi (THF)3 or Mes2SiHLi(THF)2 in pentane forms [PhBPiPr3]FeSiPh3 and [PhBPiPr3]FeSiHMes2, respectively. The complex [PhBPiPr3] FeSiHMes2 reacts with xylyl isocyanide to form [PhBPiPr3]Fe(CNXyl)2; the fate of the -SiHMes2 ligand in this reaction has not been conclusively determined. 147 The cyclohexylmethyl-substituted analog of the [PhBPPh3] ligand, [PhBPCHCy 2 3], was developed by the Peters group in 2006. CHCy CH Cy CH Cy The iron complex [PhBP 2 3]FeCl was synthesized via salt metathesis of Tl[PhBP 2 3] and FeCl2. [PhBP 2 3]FeCl was reacted with Li(dbabh) to generate the nitride [PhBPCH2 Cy3]Fe^N, in analogy to 53. Reducing [PhBPCH2 Cy3]FeCl with Na/Hg under N2 results in a [PhBPCH2 Cy3]Fe(I) species, whose structure has not been resolved, but which reacts with PMe3 and 1-adamantyl azide to form [PhBPCH2 Cy3]Fe(PMe3) and [PhBPCH2 Cy3]Fe^NAd, respectively.148 Reacting the [PhBPCH2 Cy3]Fe(I) species with 1 atm of CO2 generates {[PhBPCH2 Cy3]Fe}2(m-CO)(m-O) as the major product, and the minor products {[PhBPCH2 Cy3]Fe}2(m-Z2:Z2-oxalato) and {[PhBPCH2 Cy3]Fe(CO)}2(m-Z2:Z2-oxalato).

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

391

In 2012, the tris(di-meta-terphenylphosphino)borate ligand, [PhBPmter3]−, was introduced by the Peters group and used to synthesize the complex [PhBPmter3]FeCl by reacting [PhBPmter3]Tl with FeCl2.139 Complex [PhBPmter3]FeCl served as a precursor to [PhBPmter3]FeMe when it was treated with excess Me2Mg. Complex [PhBPmter3]FeMe was reacted with 1 equivalent of hydrazine to generate [PhBPmter3]Fe(Z2-N2H3) and methane. [PhBPmter3]Fe(Z2-N2H3) was reacted with 1 equivalent of N2H4 or NH3 to produce [PhBPmter3]Fe(Z2-N2H3)(Z2-N2H4) or [PhBPmter3]Fe(Z2-N2H3)(NH3), respectively. Both species formed [PhBPmter3] Fe(OAc)(NH3) when oxidized with Pb(OAc)4. If 1 equivalent of NH2NMe2 is added to a benzene solution of [PhBPmter3]FeMe, then the complex [PhBPmter3]Fe(Z2-NHNMe2) was obtained. [PhBPmter3]FeMe was also treated with 1 equivalent of AcOH, followed by addition of 1 equivalent of hydrazine to form [PhBPmter3]Fe(OAc)(Z1-N2H4). The anionic [PhBPPh3] and [PhBPiPr3] ligands were also used to support divalent, monovalent, and zerovalent nickel complexes (Fig. 43).149 For example, reacting 31 with (Ph3P)2NiCl2 or NiI2 generated [PhBPPh3]NiCl and [PhBPPh3]NiI, respectively. Complex [PhBPPh3]NiCl 54 was a useful precursor for several complexes including [PhBPPh3]Ni(N3) 55, [PhBPPh3]Ni(OSiPh3) 56, [PhBPPh3] Ni(O-p-tBu-Ph) 57, and [PhBPPh3]Ni(S-p-tBu-Ph) 58 (Fig. 43). Complex 54 reacted with oxygen to form [PhB(CH2P(O) Ph2)2(CH2PPh2)]NiCl or {[PhB(CH2P(O)Ph2)2(CH2PPh2)]NiCl}2, depending on the temperature and solvent. Complex 54 can be chemically reduced by Na/Hg to form [PhBPPh3]Ni(2-CH2PPh2), in which the ligand had degraded. However, chemical reduction of 54 in the presence PPh3 or tBuCN generated [PhBPPh3]Ni(PPh3) and [PhBPPh3]Ni(CNtBu), respectively. Excess NO with 54 or stoichiometric NO with [PhBPPh3]Ni(PPh3) generated [PhBPPh3]Ni(NO). Complex 31 also reacted with (Ph3P)2Ni(CO)2 in the presence of nBu4NBr to give (Ph3P)2Ni(CO)2[nBu4N]. The analogous isopropyl complex [PhBPiPr3]NiCl was prepared by the transmetalation of 34 with (DME)NiCl2. [PhBPiPr3]NiCl was generated by reacting Ni(PMe3)4 with 34. Infrared studies showed an increase in electron-releasing character for the [PhBPiPr3] ligand than in the [PhBPPh3] analog. Chemical reduction of [PhBPiPr3]NiCl with Na/Hg in the presence of excess of PMe3 formed [PhBPiPr3]Ni(PMe3). If excess tBuNC is added to [PhBPiPr3]NiCl, then the complex [k2-PhBPiPr3]Ni(Cl)(CNtBu) was formed. [k2-PhBPiPr3]Ni(Cl)(CNtBu) was reduced by Na/Hg to form [PhBPiPr3]Ni(CNtBu). [PhBPiPr3]Ni(CNtBu) was also prepared by the addition of excess CNtBu to [PhBPiPr3]Ni(PMe3). Complex [PhBPiPr3]NiCl was reacted with Li(dbabh) to generate [k2-PhBPiPr3]Ni(dbabh). 34 reacted with (Ph3P)2Ni(CO)2 in the presence of [ASN]Br to give (Ph3PiPr)2Ni(CO)2[ASN]. The ligand in 34 has also been used to form Mn(II) halide complexes, [PhBPiPr3]MnX (X ¼ Cl or I) when 34 is reacted with MnX2 salts.150 [PhBPiPr3]MnI is a precursor to the complexes [PhBPiPr3]Mn(N3), [PhBPiPr3]Mn(CH2Ph), [PhBPiPr3]Mn(Me), [PhBPiPr3] Mn(NH(2,6-iPr2-C6H3)), [PhBPiPr3]Mn(dbabh), and [PhBPiPr3]Mn(1-Ph(isoindolate), which are uncommon low-coordinate Mn complexes. Reducing [PhBPiPr3]MnI with sodium naphthalenide, followed by addition of excess CNtBu, induced the formation of a Mn(I) complex, [PhBPiPr3]Mn(CNtBu)3. Reacting Tl[PhBPiPr3] with MnBr(CO)5 generates the TldMn adduct, [PhBPiPr3] TldMnBr(CO)4. In 2018, the complexes [PhBPR3]MX (R ¼ Ph, iPr; M ¼ Ni, Co, Fe; X ¼ halide) were used for generating first-row metal silylene complexes.151 The complexes [PhBPPh3]NiCl and [PhBPiPr3]CoCl were reported to react with (THF)2LiSiHMes2 to form [PhBPPh3] Ni(m-H)(SiMes2) and [PhBPiPr3]Co(m-H)(SiMes2), respectively. Reacting [PhBPiPr3]FeBr with (THF)2LiSiHMes2 generated [PhBPiPr3] Fe(CH2–2-(SiH2Mes)-3,5-Me2C6H2). The [PhBPPh3]Ni(Z2-Bn) was found to provide a route to silylene complexes from primary and secondary silanes. For example, heating a toluene solution of [PhBPPh3]Ni(Z2-Bn) in presence of CySiH3 resulted in the formation of {[PhBPPh3]Ni(m-SiHCy)}2. In the presence of DMAP, the silylene complex [PhBPPh3]Ni(m-H)(SiHCy(DMAP)] was formed exclusively. The bis(phosphino)borate, [Ph2BPPh2] ligand was also synthesized and implemented to support the anionic platinum(II) alkyl complexes [Ph2BPPh2]Pt(Me)2[ASN] and [Ph2BPPh2]Pt(Me)(Ph)][ASN] 59 (ASN ¼ 5-azonia-spiro[4.4]nonane) (Fig. 44).152 Complex 59 was protonated in a THF solution of [iPr2EtNH][BPh4] to produce zwitterionic [Ph2BPPh2]Pt(Me)(THF) 60. The zwitterionic 60 was found to be more electron rich than its cationic relatives [(Ph2SiP2)Pt(Me)(THF)][B(C6F5)4] and [(dppp)PtMe(THF)] [B(C6F5)4], as well as undergoing benzene metalation at a faster rate. Extended thermolysis of 60 primarily generated complex molecular salt {[Ph2BPPh2]PtPh2}−{[Ph2BPPh2]Pt(THF)2}+. Interestingly, complex 60 underwent reversible [Ph2BPPh2]-metalation in a benzene solution at both the diphenylborate and the arylphosphine positions. A wide range of bis(phosphino)borates were synthesized by reacting phosphinoalkyl carbanions with the desired borane electrophile and their electronic effects examined.64 The zwitterionic palladium complex [Ph2BPPh2]Pd(THF)2][OTf] 61 was found to undergo CdH activation with trialkylamines and produced 2-coordinated iminium complexes, [Ph2BPPh2]Pd(N,C:2-NR2CHR’) (Fig. 45).153 Room temperature reactions with less polar solvents such as benzene and toluene induced formation of homoleptic Pd(I) dimer {[Ph2BPPh2]Pd}2 as the major product.

Fig. 43 The complex [PhBPPh3]NiCl as a useful precursor.

392

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

Fig. 44 Anionic platinum(II) alkyl complexes and their reactivity with acid, and benzene.

Fig. 45 Reactivity of complex 61 with trialkylamines.

Reaction of an acetonitrile solution of [ASN][Ph2BPPh2] with a benzene solution of (TMEDA)PdMe2 generated [ASN][(Ph2BPPh2)PdMe2], which was then protonated by the ammonium salt [HNiPr2Et][BPh4] to generate zwitterionic Pd(II) complex, [Ph2BPPh2]PdMe(THF).154 Complex [Ph2BPPh2]PdMe(THF) was shown to be active for CO and ethylene copolymerization, with catalytic activity comparable to its formally cationic analog,[Ph2Si(CH2PPh2)2PdMe(THF)[B(C6F5)4]. The complex [Ph2BPPh2] Pd(Me)2[ASN] will also react with [HNMe2Ph][B(C6H5)4] and acrylonitrile to form N-bound adduct [Ph2BPPh2]PdMe(NCCH] CH2).155 The complex [Ph2BPPh2]PdMe(NCCH]CH2) undergoes 2,1 acrylonitrile insertion to generate [Ph2BPPh2]Pd(CHEtCN)]n, which can then be reacted with PMe3 or pyridine to form [Ph2BPPh2]Pd(CHEtCN)(PMe3) or [Ph2BPPh2]Pd(CHEtCN)(py), respectively. In 2014 and 2015, the [Ph2BPPh2] ligand was used in the synthesis of zerovalent group VI metal complexes [Ph2BPPh2] M(CO)4[ASN] (M ¼ Cr, Mo, W) and [Ph2BPPh2]M(CO)4[NEt4], respectively, and the donor ability of the [Ph2BPPh2] ligand was explored via IR spectral analysis of these complexes.156,157 Additionally, the mononitriles [Ph2BPPh2](fac-M(CO)3(RCN)[ASN] (M ¼ Cr, R]Me; M ¼ Mo, R ¼ Et; M ¼ W, R]Et) were found synthetically useful for the introduction of sulfur dioxide and isocyanides to the metal center.156 A hybrid phosphine/pyrazolate anionic ligand, [CH2CHCH2B(CH2PPh2)(pz)2]− 62, was introduced in 2005 by Oro (Fig. 46).158 This ligand was reacted with chloro-bridged dinuclear complexes [Rh(m-Cl)(COD)]2 and [Ir(m-Cl)(COD)]2 to form new rhodium and iridium complexes [CH2]CHCH2B(CH2PPh2)(pz)2Rh(COD) and [CH2]CHCH2B(CH2PPh2)(pz)2Ir(COD) 63, respectively (Fig. 46). Lastly, this hybrid ligand framework was used in the synthesis of a polyanionic dendrimer [Li(TMEDA)]4[Si{CH2)3SiMe2(CH2)3B(CH2PPh2)(pz)2}4] to form a corresponding metallodendrimer with four rhodium atoms. The bis(phosphine)(pyrazole)borate tripodal ligands [PhBPtBu2(pz)] and [PhBPtBu2(pzMe)] 64 are two more examples of hybrid phosphine/pyrazolate anionic ligands (Fig. 47).159

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

393

Fig. 46 Synthesis of mononuclear diolefin/scorpionate complexes.

tBu Me Fig. 47 Bis(phosphino)(pyrazolyl)borate tripodal ligands [PhBPtBu 2 (pz)] and [PhBP2 (pz )].

A few related aluminates come from Bourisou et al., who in 2008160 created the first diphosphine-alane (DPA) system which contained an Al(III) center featuring two C-bound arylphosphines and a chloride. Reacting this species with ClAu(SMe2) generated the zwitterionic complex, which had undergone chloride transfer from Au to Al. Similar reactivity is seen when this DPA ligand161 is reacted with CuCl giving chloride-bridged polymeric species, which also displays Cl transfer from Cu to Al. Expanding the scope of diphosphine–alane systems, Bourissou subsequently created a trisphosphine–alane162 system (TPA). This ligand was again capable of coordinating to Au and Cu by reaction of TPA with CldAudSMe2 or CuCl to give the corresponding zwitterionic species after chloride abstraction by the Al center. In some of these systems the four-coordinate aluminate center acted as a Z-type ligand. This contrasts with ligands containing borates, as the boron center does not have low enough energy empty orbitals to interact with metal nonbonding electrons.

1.12.3.1.4

N- and S-Donor scorpionates and related bidentate systems

Poly amino and thioether borates can be synthesized via nucleophilic attack of alkyllithium reagents on substituted haloboranes, but this method requires the R groups of the thioether to lack alpha protons (with the exception of the methyl derivative).63 1.12.3.1.4.1 Coordination chemistry/reactivity While reports of catalysis and new reactivity in these systems have not been thoroughly investigated, their coordination chemistry, particularly for the thioether derivatives,63 has been well studied. An interesting report from Betley and Peters65 details the coordination chemistry of a bis(amino)borate 65 (Fig. 48). The electronics exhibited through this particular N-donor ligand, as well as similar ligands involving different donor atoms, present an interesting case study. Though clearly the charge is separated, the ensuing bis(amino)borate complex demonstrated diminished Lewis acidity, as evidenced by CO stretching frequency analysis of 66 (Fig. 48) when compared to the formally cationic metal complexes stabilized by weakly coordinating anions. Thus, it has been deemed appropriate to describe complexes utilizing these borate complexes as zwitterionic. The resulting bis(amino)borate is analogous to anionic 1,3-diketimines as well as charge neutral tertiary amine chelates such as tetramethylethylenediamine (TMEDA).

1.12.3.1.5

Scorpionates and related bidentate systems featuring NHCs

1.12.3.1.5.1 Ligand synthesis N-Heterocyclic Carbene (NHC) bearing borates 67 and 68 are an interesting class of multidentate ligands that complement analogous phosphine species (Fig. 49). The precursors to these ligands are poly(imidazolyl)borates that can be synthesized in

Fig. 48 Synthesis of a zwitterionic Rh dicarbonyl complex of a bis(amino)borate ligand.

394

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

Fig. 49 Generic tris- and -bis (NHC) borates. General formula for 67 RB(R’Im)3 and 68 R2B(R’Im)2; protonated forms will be represented by generic formulas RB(R’ImH)3 and R2B(R’ImH)2, respectively. Drawn as free carbenes for clarity. R ¼ H, F, aryl; R’ ¼ alkyl, aryl.

two main ways: (1) The imidazole subunits are installed onto B nuclei followed by alkylation to yield cationic carbene precursors. (2) The imidazole is functionalized prior to substitution onto the appropriate haloborane derivative.163 To achieve the desired metal complex, the carbenes were typically generated in solution followed by the addition of a metal halide. Alternatively, treating the precursor with Ag2O makes the Ag(NHC) complex that could be transmetallated to the desired metal complex using appropriate metal halide precursor. 1.12.3.1.5.2 Electronics Like the Tp and Bp systems, bis and tris (NHC)borates are considered 4e− donors. The electron-donating properties have been studied in tris(NHC)borates by probing the NO stretching frequency in tetracoordinate {NiNO} complexes.164 Varying the N-substituent led to small changes in NO stretching frequencies; however, a greater effect was found by modifying the heterocycle. The NHC donors in descending order of electron donating properties were imidazol-2-ylidene > benzimidazol-2-ylidene > 1.3,4-triazol-2-ylidene. Additionally, all the above NHCs were more donating than the Tp and Bp ligands. Another example of a bis(NHC)borate complex was the nickel nitrosyl compound, Ph2B(tBuIm)2Ni(NO)(PPh3).165 This compound exhibited a trigonal pyramidal structure, linear NO ligand, and long NidP distance, and is thus better described as a three-coordinate complex. The NO stretching frequency was similar to that of 1,3-bis(2,6-diisopropylphenyl)-imidazolidin-2-ylidene)Ni(NO)I and lower than that of the other known three-coordinate nickel nitrosyl complexes, which is consistent with these ligands being very strong donors. 1.12.3.1.5.3 Coordination chemistry/reactivity Since Fehlhammer first reported the tris(NHC)borate complexes in 199666 and later bis(NHC)borate complexes in 2001,166,167 this field has grown rapidly with interesting discoveries in coordination chemistry and exciting applications from catalysis to single molecular magnets. Both bis- and tris-NHC ligands tend to form homoleptic complexes when the R groups have minor steric contributions. Since these species are coordinatively saturated, their further reactivity is limited. In contrast, heteroleptic complexes can be stabilized and isolated when wingtip groups are more sterically encumbering around the metal center. The Smith group has explored the reactivity of iron complexes stabilized by these ligand scaffolds. Low valent iron carbonyl compounds were isolated from the reaction of [PhB(MesImH)3][OTf−]2 with LDA to yield [Li][PhB(MesIm)3] which when reacted with FeCl2(THF)1.5 yields the iron(II) complex PhB(MesIm)3FeIICl, 69.168 The FedCl bond in 69 could be cleaved by salt metathesis with [K+][B(C6F5)4−] in the presence of CO to yield [PhB(MesIm)3FeII(CO)3][B(C6F5)4]. One electron reduction of [PhB(MesIm)3FeII(CO)3][B(C6F5)4] with KC8 yields PhB(MesIm)3FeI(CO)2, losing [K+][B(C6F5)4−] and a CO moiety in the process. Further reduction with KC8 led to the formation of a dimeric, dianionic, Fe complex [K]2[PhB(MesIm)3Fe(CO)2]2 70, with interesting side on coordination of the CO ligands to K+ forming a dimeric species (Fig. 50). An extension of this work features the bis(NHC)borate ligand Ph2B(tBuImH)2 which, when deprotonated and reacted with FeCl2, yields the iron(II) complex Ph2B(tBuIm)2FeIICl(THF) 71 (Fig. 51).169 Reduction of 71 with KC8 in toluene yields Ph2B (tBuIm)2FeI(toluene), which when subjected to CO atmosphere forms Ph2B(tBuIm)2FeI(CO)3. Further reduction of these low coordinate iron species with KC8 leads to divergent reactivity based on whether the 2,2,2-cryptand additive was employed in the reaction. When the cryptand additive was employed, the reaction yielded a monomeric anionic iron species K(crypt)[Ph2B (tBuIm)2Fe(CO)3]. When the cryptand was not employed, a dimeric species [K]2[Ph2B(tBuIm)2Fe(CO)2]2 72 was formed by

Fig. 50 Common iron precursor 69 and dimeric iron(0) complex 70.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

395

Fig. 51 Bis(NHC) analogs of common iron precursor Ph2B(tBuIm)2FeIICl(THF) 71 and dimeric iron(0) complex [K+]2[Ph2B(tBuIm)2Fe(CO)−2 ]2 72.

coordination of the CO ligands to K+. This work demonstrated the ability of the Ph2B(tBuIm)2 ligand to stabilize low coordinate iron species. It was later discovered that the monomeric species K(crypt)[Ph2B(tBuIm)2Fe(CO)3] undergoes thermally activated CdH activation with the iron center inserting into the CdH bonds of the tert-butyl group.170 Highly oxidized iron nitrido species were also isolated by Smith et al. using both bis and tris(NHC)borates; these are interesting because of their relationship to industrial and biological nitrogen reduction processes.171 Reaction of PhB(tBuIm)3FeIICl 73 with NaN3 under UV irradiation yielded the iron(IV) nitrido species, 74 (Fig. 52).172 These iron(IV) nitrido species react with radicals,173 alkenes,174–176 and nucleophiles (e.g., phosphines,177 isocyanates,178 carbon monoxide,178 silanes,179 cyclopropylidenes180) and many other derivatives.181,182 One electron oxidation reactions employing a ferrocenium cation oxidant led to the formation of a formally Fe(V) species [PhB(tBuIm)3FeVN][BArF24] 75 (Fig. 52).183,184 Compound 75 produced ammonia in 89% yield in the presence of 3 equivalents of cobaltocene and 15 equivalents of water. Furthermore, the iron(IV) nitrido species 74 engages in partial nitrogen atom transfer reactions to create heterobimetallic complexes by reacting with V(Mes)3 to yield PhB(MesIm)3FeIIN] V(Mes)3.182 Reacting 75 with molybdenum phosphide yielded a nitride.185 Alternatively, a cyanide ligand was substituted onto PhB(tBuIm)3FeIIF with TMS-CN to create PhB(tBuIm)3FeIICN, in which the Lewis basic nitrogen on the cyano group can bind to another metal demonstrated by coupling with a Mo(III) precursor, enabling high magnetic exchange coupling through cyanide.186 Several years after this report, bis(NHC)borates were also employed to undergo similar chemistry. The compound H2B (MesIm)2FeIII(THF) 76 was synthesized by reacting 2 equivalents LDA with [H2B(MesImH)2]I followed by addition of FeCl2(THF)1.5.187 Iron complex 76 was then subjected to a one-electron reduction with KC8 in the presence of

Fig. 52 Synthetic route toward highly oxidized iron (V) nitride species 75.

Fig. 53 Synthetic route toward highly oxidized iron (VI) bis imido species 80.

396

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

divinyltetramethylsiloxane (dvtms) to yield H2B(MesIm)2FeI(dvtms) 77 (Fig. 53). Reaction of 77 with 2 equivalents MesN3 yielded bis(imido) compound H2B(MesIm)2FeV(]NMes)2 78. When subjected to one electron oxidation using the ferrocenium cation, the cationic species{[H2B(MesIm)2FeVI(]NMes)2][BR4] (R ¼ F and Ph) 79 was formed. This compound was stable in the solid state but slowly decomposed in solution. These highly oxidized iron compounds adopt unusual seesaw geometries and the covalent bonding interactions in Fe(V) and Fe(VI) were thoroughly investigated.187 Interestingly, a dimeric iron hydride complex has been isolated by reacting Ph2B(tBuIm)2FeIICl(THF), akin to 76, with NaHBEt3. This [Ph2B(tBuIm)2FeIIH]2 dimer contains two geometrically distinct iron centers as one is tetrahedral and the other square planar.188 Both the iron centers are doubly oxidized but exhibit different spin states. However, no such hydride complex was observed when the tris(NHC)borate iron complexes were subjected to similar conditions. The reaction of PhB(MesIm)3FeIICl with NaHBEt3 yielded PhB(MesIm)3FeII(HBEt3), where the triethylborohydride ligand was bound to Fe by an FedHdB bond and two CdH agostic interactions.168 PhB(MesIm)3FeII(HBEt3) reacted with CO to yield the iron hydride PhB(MesIm)3FeII(CO)2H discussed above.168 Alkene isomerization catalysts have also been developed using iron-NHC borate scaffolds. Reaction of Ph2B(tBuIm)2FeIICl(THF) and LiCHt2Bu in pentanes yielded Ph2B(tBuIm)2FeII(CHt2Bu).189 Subsequent reduction with KC8 in the presence of 2,2,2-cryptand under N2 atmosphere led to the formation of Ph2B(tBuIm)2Fe(CHt2Bu)(N2). Addition of 1-hexene to Ph2B(tBuIm)2Fe(CHt2Bu)(N2) displaced the labile dinitrogen and isomerized completely to 2-hexene in a 4:1 trans:cis ratio.189 While there have been numerous contributions to bis and tris(NHC)borate iron chemistry, there has been substantially less progress made in exploring other transition metals. Heteroleptic manganese(I) carbonyl compounds have been synthesized using tris(NHC)borate scaffolds.190 These compounds readily react with oxygen to form dicationic, Mn(IV) homoleptic tris(NHC) borates190 that were employed in photoluminescence experiments.191 Further, they were used in the synthesis of homoleptic Mn(III) compounds via one electron reduction and compared to tris(pyrazolyl)borate analogs.192,193 Mn(IV) nitrides have also been prepared with chemistry analogous to 74 and investigated as single molecular magnets.194 Cobalt complexes featuring tris(NHC)borate scaffolds have also been synthesized. PhB(tBuIm)3CoCl has been demonstrated to undergo reversible protonation/deprotonation and can be reacted with MeLi or MeMgBr to form PhB(tBuIm)3CoMe.195 Also, Co(III) imido complexes have been synthesized through hydrogen atom abstraction of PhB(tBuIm)3Co(NHtBu).196 Cobalt azido complexes have been synthesized by reacting PhB(tBuIm)3CoCl with NaN3.191 Attempts to transform the resulting PhB(tBuIm)3CoN3 complex by photolysis and thermolysis were unsuccessful. Nickel complexes have been mainly explored to assess the electron-donating properties of these ligands. Of note is a report of 4-coordinate d8 nickel complexes, which were synthesized by reacting [Ph2B(tBuIm)2Li(Et2O)2] with NiCl2(PMe3)2 to yield Ph2B(tBuIm)2Ni(PMe3)Cl.165 Subsequent reaction with PhCH2MgBr gives the square planar Ni(II) d8 complex Ph2B(tBuIm)2Ni(Z3-CH2Ph). Homoleptic complexes featuring group 11 metals, ruthenium and osmium have also been synthesized and previously reviewed by Santini, Marinelli, and Pellei.163

1.12.3.1.6

Mixed NHC borates

Sadow et al. reported the tridentate borate ligands bearing two 4,4-dimethyl-2-oxazoline (Ox) moieties and one substituted imidazole moiety. These can be introduced to metal centers via deprotonation/transmetalation procedures yielding an k3-N,N,C chelate via two oxazolinyl nitrogens and the NHC carbon. Deprotonation of the borate ligand was performed with KCH2Ph197,198 and dialkylzinc199 reagents, both providing useful nucleophilic precursors. Sadow first introduced these ligands onto rhodium and iridium compounds that have been previously reviewed.163,198 Partial deoxygenation catalysts have been developed by reacting the carbenoid precursor, [PhB(OxMe2)2(MesIm)−] 80, with ½ equivalent of [Rh(m-Cl)(CO)2]2 to yield

Fig. 54 Synthetic route toward partial deoxygenation catalyst 83 and an example of a catalytic reaction demonstrating the mixture of products obtained.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

397

PhB(OxMe2)2(MesIm)Rh(CO)2 81 (Fig. 54).197 Subsequent treatment of 81 with PhSiH3 in benzene caused the loss of one CO ligand and the oxidative addition of the SidH bond creating PhB(OxMe2)2(MesIm)RhH(CO)(SiH2Ph) 82. Treatment with Lewis acid B(C6F5)3 removes the metal hydride and causes a second oxidative addition of an SidH bond, followed by insertion of the silane moiety into the RhdN bond of one of the oxazolines to yield 83. 1.12.3.1.6.1 Scorpionates and related bidentate systems with imidazole-2-thione borates 1.12.3.1.6.1.1 Ligand synthesis In general, these ligands are prepared by the reaction of mono-N-substituted imidazole-2-thione derivatives with an appropriate borohydride. Variations in temperature and solvent enable the isolation of mono, di-, and tri-substituted imidazole-2-thione borate ligands (Fig. 55).200 The R substituents on the boron backbone tend to be H but other aryl groups have also been explored. The N-R’ moiety has been replaced by other atoms, such as sulfur, but there are no reports of catalytic activity in which these derivatives perform better than the imidazole-2-thione derivatives.201 1.12.3.1.6.1.2 Coordination chemistry/reactivity Imidazole-2-thione donor borates coordinate in a variety of different modes and have multiple degradation pathways, which have been previously reviewed. So far, there have been only two examples of catalytic activity with these ligands. Coupling reactions have been demonstrated with Ag(I) complexes bearing a DIPP-NHC ligand as well as a bis(imidazole-2-thione)borate ligand.201 Also, catalytic oxygen atom transfer has been observed in Mo(VI) oxo compounds bearing tris(imidazole-2-thione)borate ligand.202 Treatment of (DIPP-NHC)AgCl with the Na salt of 85 (R ¼ H; R’¼Me) in chloroform precipitated NaCl and yielded complex 86 (Fig. 56).201 Complex 86 was catalytically active in A3 coupling reactions of benzaldehyde, piperidine, and phenylacetylene in 10:1 H2O/THF solvent system demonstrating the water stability of the borate backbone and heterocycles. Furthermore, the Ag(I) complexes had comparable performance to a DIPP-NHC supported gold(I) SMe2 complex demonstrating the possibility of using cheaper metals for similar catalysis. The reaction of MoO2Cl2 with Na+[85 (R]H; R’]Me)] in methanol led to complex 87 in good yields (Fig. 56).202 Complex 87 was capable of catalyzing the transfer of oxygen from DMSO to PPh3, generating O]PPh3 in acetonitrile. 1.12.3.1.6.2 Ansa-bridged metallocenes While not a scorpionate ligand, boron-bridged ansa-metallocenes are relevant to this section as they are isolobal to homoleptic scorpionate complexes, and have a charge that is separated from the metal. While they will not be discussed in detail herein, it should be noted that some metallocenes feature distal borates appended to Cp∗ ligands via one of the methyl groups.203–206 With respect to the former ansa-bridged variants of these borates, these species were highly sought after for a time as cationic metallocenes akin to 88 (Fig. 57) and are believed to be the active species in olefin polymerization catalysts.207 However, these species could not be isolated in similar carbon and silicon analogs. Thus, it was postulated that a boron-bridged metallocene species, 88, could

Fig. 55 Generic structure for tris- (84) and bis-(imidazole-2-thione) borates (85); R ¼ H, aryl; R’ ¼ t-Bu, aryl.

Fig. 56 Poly(imidazole-2-thione)borate derived catalysts 86 and 87.

398

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

Fig. 57 Sought after active species of olefin polymerization 88 and typical ligand architecture borato ansa-metallocenes 89 [M ¼ Ti(IV) or Zr(IV)].

stabilize this coordination environment due to the close electrostatic interactions. Species like 89 (Fig. 57) have been the main platforms for this chemistry, where there is a Lewis basic L-type ligand at the boron center.207 1.12.3.1.6.2.1 Ligand synthesis These ligands are generally prepared by either deprotonating Cp rings attached to the boron nuclei or by activation of pendant groups off the Cp ring.207 Precursors for the latter species have been prepared by the reaction of (1,1-(Me3Si)(Me3Sn)-cyclopentadiene) 90, with PhBCl2 to furnish (Me3SiCp)2BPh 91 (Fig. 58).208 To further control the ligand environment, one can perform nucleophilic substitution on (Me3SiCp)2BBr 92 as reported by Shapiro, which aids in diversifying the borate backbone.207 Different backbone structures have been investigated to which the nature of the B-L adduct (89, Fig. 57) has been found to be very important. The adduct needs to be stable enough to remain unreactive with strong alkyl nucleophiles used to alkylate the metal centers in preparation of olefin polymerization catalysts. The SMe2 adduct initially synthesized was found to be too labile for these subsequent reactions.209 Ligand substitution reactions were able to be performed with phosphines209 and phosphorus ylides,210 which eventually led to catalytic activity.210 Derivatives using indene rings instead of cyclopentadiene have also been synthesized by reacting lithiated indene rings with appropriate haloborane.210,211 These rings could be deprotonated again in the preparation for transmetallation onto other metals. Bochmann later synthesized derivatives with no ylidic adduct present using alkyl and aryl nucleophilic substitutions beginning with species like 93.212 1.12.3.1.6.2.2 Coordination chemistry/reactivity Shapiro reported that the treatment of (Me3SiCp)2BPh with ZrCl4(SMe2) caused the loss of Me3SiCl to furnish ansa-bridged zirconocene with a tetracoordinate borate backbone 94, where R]Ph, L ¼ SMe2.209 In the course of the reaction, however, the boron forms an adduct with SMe2 in solution causing the formation of an ylide. However, some complexes contacting ylidic adducts have been observed to function as competent olefin polymerization catalysts competitive with the more common carbon and silicon-bridged metallocenes.210 The SMe2 on the boron backbone has been substituted via ligand exchange reactions with phosphines209 and phosphorus ylides210 and both were demonstrated to be stable for alkylation reactions and were competent olefin polymerization cocatalysts. Bochmann later demonstrated the utility of nucleophilic substitutions of alkyl and aryl lithium or Grignard reagents in selectively substituting ZrdCl over BdCl bonds (Fig. 59).212 It was critical to begin with (Me3SiCp)2BCl 93, instead of Shapiro’s (Me3SiCp)2BPh, 91, to simplify subsequent substitutions on the boron center. Compound 93 reacted with ZrCl4(SMe2)2 to yield (SMe2)ClBCp2ZrCl2 94. Two equivalents of C6F5Li or C6F5MgBr selectively alkylated ZrdCl over BdCl bonds yielding 95, while an excess beyond 4 equivalents yielded 96. Reaction of 96 with 2 equivalents MeLi alkylated at the zirconium center, yielding ethylene polymerization cocatalyst 97. The synthesis leading up to 97 is low yielding and evidence of decomposition was found. Typical methide abstracting reagents were necessary for catalytic activity to be observed and species 88 was never isolated. Titanocene and zirconocene derivatives containing borato-bridged ansa-metallocenes have been previously reviewed207 as well as new avenues toward boron-bridged ferrocenes.213

Fig. 58 Synthesis of common boron-bridged ansa-metallocenes ligand precursors 90–93.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

399

Fig. 59 Bochmann’s synthesis of anionic borato-bridged ansa-metallocenes 96 and 97.

Fig. 60 H+ complex 98 and its utility in the synthesis of hydroamination precatalysts 99 and 100.

1.12.3.1.7

Mixed Cp systems

Sadow et al. introduced mixed Cp systems on scorpionate-type scaffolds containing two 4,4-dimethyloxazoline groups, akin to mixed NHC borate ligand 80. These ligands bind in an L3X2 fashion. A convenient means toward metal complexes with this ligand utilizes the acid, 98, which reacts Zr(NMe2)4 or Hf(NMe2)4 (Fig. 60).214 The ensuing complexes 99 and 100 were investigated as active precatalysts for hydroamination.214–218 Lanthanum, yttrium, and neodymium aluminate complexes were synthesized from 98 and the corresponding M(AlMe4)3 salts (M]Y, La, Nd). These compounds were investigated for their catalytic activity in the alumination of terminal alkyne substrates.219,220

1.12.3.2

Diimines and related systems

Diimine ligands have been extensively investigated as they have demonstrated great utility in olefin polymerizations incorporating polar monomers initially reported by Brookhart. Investigations on improving, understanding, and modifying catalytic activity were undertaken. By taking Brookhart’s catalyst as a starting point and modifying substituents around the diimine ligand, calculations were able to deliver insight into possible modifications that could aid in improving catalyst performance.47–49 An initial study determined that bulky ortho-aryl substituents and incorporation of anionic character in the ligand framework both aided in stabilizing p-type complexation of the alkene in polar monomers over s-type interactions, which lead to catalyst poisoning.49 Further computational experiments by Ziegler and coworkers investigated the energy barriers associated with monomer insertion and other side reactions. Specifically, different anionic (BF3−, SO3−, and BH3−) and electron-withdrawing substituents were substituted in R group positions (Fig. 61) and their ensuing conformations were probed. From the computational data, it was postulated that incorporation of BF3− in the ortho positions of the diimine ligand would aid in suppressing observed oligomerization reactions. Furthermore, many of the BF3−-substituted derivatives had the highest propagation rates. Bazan and coworkers have developed the approach of utilizing a-iminocarboxamidato and a-iminoenamido ligand platforms. He initially synthesized a-iminocarboxamidato ligand platforms via deprotonation of the amide 101 with KH, followed by complexation in a bidentate fashion (Fig. 62).221,222 This NiII species, 102, reacts with 2 equivalents of B(C6F5)3 to form activated complex 103, where the Lewis acid binds to the carbonyl fragment on the amide ligand to withdraw electron density. Activation of

400

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

Fig. 61 Ziegler’s depiction of structural modifications made in his computational experiments.

Fig. 62 Bazan’s synthesis of a-iminocarboxamidato NiII complexes and their activation via treatment with B(C6F5)3.

complex 103 with trialkylaluminum reagents has also been reported.223 Of note, bulky aryl substituents caused O-chelation in initial binding to the NiII center but were isomerized to complexes of type 103 when subjected to a Lewis acid.222 The complexes were found suitable for the production of high-molecular-weight polyethylene.222,224 O-chelation and isomerization to C-chelation were also observed for enolate derivatives of 102, which replace the amide functionality for an acetate group.225 Pyridinecarboxamidato NiII complexes, containing a 2-pyridine moiety in place of the imine functionality in 102, were synthesized and investigated for similar reactivity.226 Following these findings, the a-iminoenamido ligands were synthesized as isoelectronic derivatives of a-iminocarboxamidato ligands discussed above. Generation of the enamine from reaction of KH with 104 followed by addition of [Ni(methallyl) Cl]2 yielded complex 105 (Fig. 63).227 Subsequent treatment of 105 with a Lewis acid yielded compounds of type 106. These compounds were investigated utilizing B(C6F5)3 and Al(C6F5)3 as Lewis acids, and replacing Z3-methallyl for Z3-benzyl fragments as well.227 Zwitterionic complexes of type 106 were active toward ethylene polymerization but addition of THF to solutions of the catalyst demonstrated reversion to complex 105. These results were used to explain the importance of excess Lewis acid in polymerization activity as a scrubber for coordinating molecules that lead to catalyst deactivation and eventually poisoning.

1.12.3.2.1

b-Diketiminate platforms and related systems

Another important subclass of ligands in this field are the b-diketiminate (“nacnac”) ligand derivatives 107 and 108 (Fig. 64). Generally, the identities of the R group in these ligand sets have a p-system embedded in order to create an avenue for withdrawal of electron density of the metal center. Lewis acid activation through the available p-system of ligands 96 and 97 is discussed herein.

Fig. 63 Bazan’s synthesis of a-iminocarboxamidato NiII complexes and their activation via treatment with B(C6F5)3.

Fig. 64 Modified NacNac ligand platforms and their Schiff base derivatives. The R substituent is the exocyclic position where the Lewis acid binding site is located.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

401

Fig. 65 Ligand isomerization and preparation of catalytically active species 110 by treatment with B(C6F5)3.

Initially Bazan227 et al. investigated NiII complexes bearing derivatives of 107 for olefin polymerization. Reaction of potassium salt derivatives of 107 with NiII(Z3-benzyl)PMe3Cl led to an O,O-chelated NiII complex, 109.228 Subsequent treatment of 109 with B(C6F5)3 binds to one of the chelating oxygen atoms leading to ligand isomerization (Fig. 65). Another equivalent of B(C6F5)3 removes the coordinated phosphine accessing the catalytically active species 110. The active species was capable of producing oligomeric species in the presence of ethylene.

1.12.3.2.2

Miscellaneous ligand platforms

Evidence for improving polymerization performance through the incorporation of anionic ligands has also been demonstrated in other ligand platforms. The initial investigations of Bazan et al. into zwitterionic NiII complexes capable of tandem ethylene polymerization utilizing this methodology began with modifying SHOP catalyst derivatives popularized by Keim.229 NiII Z3methallyl complex of ortho diphenylphosphine benzoic acid, 111, was initially studied by Keim, who investigated neutral and cationic derivatives for olefin polymerization activity (Fig. 66).229 Treatment of o-diphenylphosphinebenzoic acid with NiII(Z3-methallyl)2 yielded (o-diphenylphosphinebenzoate)NiII(Z3-methallyl), 111. Treating complex 111 with B(C6F5)3 led to the formation of a BdO bond between the carboxylate and borane yielding zwitterionic complex 112 (Fig. 66).230 Both 111 and 112 are reactive toward ethylene but zwitterionic complex 112 had a higher activity and yielded higher MW polymers. However, it was discovered that the methallyl functionality was less reactive than the propagating species causing only a fraction of the nickel centers being relevant in catalysis. This led to the synthesis and investigation of isoelectronic Z3-benzyl fragment replacing the less reactive Z3-methallyl fragment. Treatment of ortho-diphenylphosphine sodium benzoate with benzyl chloride and Ni(COD)2 led to the formation of a NiII dimer bridging through the carboxylate functionality, 113 (Fig. 67).231 Compound 113 could be treated with 2 equivalents of B(C6F5)3 to yield monomeric NiII species

Fig. 66 Keim’s (o-diphenylphosphinebenzoate)NiII(Z3-methallyl) complex, 111, and Bazan’s “activated” complex 112.

Fig. 67 Bazan’s synthesis of NiII complexes 113 and 114 bearing Z3-bezyl fragments.

402

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

114, which was more active than the corresponding methallyl fragment 112. Jordan et al. demonstrated the efficacy of this approach for other ortho-substituted ligand frameworks, and demonstrated Lewis acid binding onto ortho-substituted diarylphosphine -sulfonates, −amidos, −phosphonates, −phosphinates, as well as other non-phosphine-related donors.26,232–236 Binding of Lewis acid to a carbonyl substituent of the ligand can withdraw more electron density from the metal leading to higher activity. The higher activity is not always accompanied by higher molecular weight since corresponding increases to chain transfer constants can lead to lower molecular weight distributions, which is the case for ortho-substituted phosphine–sulfonate PdII complexes.26

1.12.3.3

Monodentate ligands with proximal and distal borate ligand substituents

Although underdeveloped compared to their multidentate analogs, monodentate ligands featuring distal and proximal four-coordinate borates are known. In this section we will highlight some of the most interesting systems. Monodentate Phosphines Featuring Distal Borates: One of the earliest reports of ligands in this class is from Peters and coworkers.237 They synthesized a series of monodentate phosphines 115–116 appended with the tetraphenylborate anion (Fig. 68). All four ligands reacted with [(NBD)RhCl]2 to form [NR4][(Ph3BPR2)RhCl(NBD)] salts. The [NnBu4][(Ph3BPp-iPr2)RhCl(NBD)] salt reacted with TlPF6 to generate a zwitterionic rhodium species. The ammonium salts 115–116 reacted with (COD)PtMe2 to afford disubstituted products [NR4]{[Ph3BPR2]2PtMe2} exclusively as the cis isomers. Stoichiometric addition of B(C6F5)3 to [NR4] [(Ph3BPR2)2PtMe2] generated [NR4][trans-(Ph3BPR2)2PtMe(solv)] and [NR4][Me(B(C6F5)3)]. Lastly, the [NR4][Ph3BPR2] ligands were used to promote Suzuki cross-coupling reactions (Fig. 69). Subsequently, Spivak and coworkers238 synthesized a series of anionic half-sandwich ruthenium–arene complexes 117 containing (phosphino)tetraphenylborate ligands (Fig. 70). Treating 117 with halide-abstracting reagents like AgNO3 or AgOTf generated the zwitterionic species [RuCl(L)(Z6-p-cymene)(PR2(p-Ph3BC6H4))] (L ¼ pyridine or MeCN). However, no catalytic applications have been reported for these complexes.

Fig. 68 Synthesis of the anionic monodentate phosphino borate ligand [Ph2,B(CH2PPh2)2]− (abbreviated [Ph3B].

Fig. 69 Suzuki cross-coupling between phenylboronic acid and p-chlorotoluene, p-chloroacetophenone, or 1,4-dichlorobenzene using the anionic [Ph3BPm-iPr2] ligand.

Fig. 70 Synthesis of anionic half-sandwich ruthenium-arene complexes containing (phosphino)triphenyl borate ligands.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

403

More recently, Stephan et al.239 reported a series of PdH oxidative addition reactions with phosphonium ions containing pendant fluorinated aryl borate substituents to access zwitterionic Pt(II) compounds. In the same publication, they disclosed the synthesis of a phosphine containing a long-chain alkyl linker attached to a fluorinated aryl borate and demonstrated its direct coordination to Ni(II) and Ru(II). Although the coordination chemistry of these zwitterions was not thoroughly discussed, reactivity of the ensuing complexes was investigated. Systems featuring proximal borate moieties: Formed indirectly via PdH oxidative addition, some of the first examples of phosphine complexes appended with a proximal borate anion are from Manners et al.240 (Fig. 71). They reported a series of Pt complexes of phosphine–borate anions of the form [RPhPBH3]− and structurally characterized them but reported minimal reactivity studies. Similar ligand scaffolds have been prepared from phosphine boranes with a hydrocarbon linker which, when exposed to transition metal fragment, abstracts chloride/methide and generates a phosphine bound to a metal and a pendant borate moiety. For example, Bourissou et al.241 utilized the latter approach in the reaction of iPr2PC6H4BCy2 with [(allyl)PdCl]2 to give Pd − P and B − Cl bonds. Likewise, Tilley et al.242 reacted phosphinoethylboranes with Ni–methyl complexes to give zwitterionic phosphinemethylborate complexes. Another unique class of monodentate phosphine borates are triptycene derivatives.243 Subsequently, Piers244 and Jordan245 introduced aryl phosphine ligands ortho-functionalized with a proximal BF3− substituent independently. They investigated the coordination chemistry of these ligands and their ability to mediate olefin polymerization (Fig. 72). In nonpolar media, the BF3− group interacts functionally with the metal center. Furthermore, while not competent as an olefin polymerization catalyst, these complexes readily dimerize ethylene. Another interesting and unique system with a proximal borate moiety is diphenylphosphidoboratabenzene, [Ph2PBC5H5−] 118,246,247 which is an anionic isosteric analog of the ubiquitous PPh3 (Fig. 73). The borate differs in these systems as it is tricoordinate, planar, and is involved in a 6-membered aromatic ring. They showed that this class of ligands can bind transition metals as exemplified by the formation of the zwiterrionic Rh(I) compound 119 and the Fe(II) complex 120 (Fig. 73).

1.12.3.3.1

Mono-NHC borates

Monodentate NHC’s Featuring Distal Borates: While investigating carbene–borane Frustrated Lewis Pair (FLP) systems, Tamm et al. discovered that hydrogen splitting activity was no longer observable after 2 h in 1,3-di-tert-butylimidazolylidene-B(C6F5)3 systems without H2 atmosphere. This was due to the system quenching itself producing zwitterionic imidazolium 122 (Fig. 74).248 Compound 121 did not exhibit any FLP activity due to the much stronger adduct formed in the Cbackbone–B bond versus the Ccarbene–B interaction. Tamm et al. sought the deprotonated form of 121 as they wanted to investigate organometallic complexes bearing this NHC, unfortunately all attempts were unsuccessful.249 Formation of NHCs bearing weakly coordinating moieties on the backbone carbon was not possible until pioneering work from the Robinson group, who discovered that bis-N,N-substituted NHCs deprotonated with n-BuLi react selectively at the backbone carbon with electrophile.250 Adapting these synthetic protocols,

Fig. 71 Manners synthesis of [RPhPBH−3 ] via PdH oxidative addition.

Fig. 72 Piers and Jordan’s dimerization catalysts.

Fig. 73 Fu’s anionic analog of triphenyl phosphine.

404

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

Fig. 74 Tamm’s synthesis of monodentate NHC borates adapted from Robinson et al.

Tamm et al. synthesized LidNHC complexes bearing anionic borate substituents, 122–125, by reacting the doubly deprotonated imidazoliums with one equivalent of B(C6F5)3. The utility of this class of ligands has been demonstrated via numerous examples of transmetalation reactions that have yielded useful catalysts. Au(PMe3) and Au(THT) complexes were initially reported along with the ligand synthesis by Tamm et al. via reactions of 122 and 123 with ClAuPPh3 and ClAu(THT), respectively.249 Reactions with ClAuPPh3 afforded AuPPh3 complexes in good yield following precipitation of LiCl. The reaction of 122 with ClAu(THT) led to the formation of metallic Au, but the reaction of 123 led to the desired Au(THT) complex, 126 (Fig. 75). Complex 126 was subjected to a model enyne rearrangement (Fig. 75, right) and compared to the best performing cationic catalyst 127 known at the time (Fig. 75). Zwitterionic complex 126 demonstrated comparable performance to 127, demonstrating its utility in enyne rearrangements.249 Following their work with Au(I) complexes, Tamm et al. focused their attention on Rh(I) and Ir(I) complexes of ligands 122–124 as well as ligands with modified borate substituents (i.e., B(p-Tol)3 and B(m-XyF6)3) on the NHC backbone.251 Reactions of 122–124 and their modified borate analogs with [M(COD)Cl]2 (M ¼ Rh or Ir) yielded complexes of type (NHC)M(COD). The iridium analogs were exploited as hydrogenation catalysts, among which the Ir(COD) bearing 124 as the NHC, 129 (Fig. 76) was found to be the most effective. This family of zwitterionic complexes was capable of performing hydrogenation reactions in nonpolar solvents (i.e., hexane and cyclohexane), in which typical cationic catalysts have difficulty. Neat hydrogenations were performed with catalyst loadings as low as 10 ppm, under 8 atm of H2. Furthermore, directed hydrogenation of rapeseed oil could be performed with a very high degree of stereospecificity (product ratio > 99:1) by 128 (Fig. 76).251 Compound 128 and its N-mesityl derivatives were subsequently modified via incorporation of an L-type ligand onto the Ir center, 129 (Fig. 76).252 Compounds 128, 129, and the N-mesityl derivatives of 128 act as Hydrogen Isotope Exchange (HIE) catalysts at 5 mol% loading, 1 atm D2 and differing activity depending on the solvent used. 129 and its N-mesityl derivatives were also capable of performing reactions in nonpolar solvents.252

Fig. 75 Tamm’s zwitterionic catalyst 126, highly active catalyst 127, and the model enyne cyclization reaction.

Fig. 76 Hydrogenation catalyst 128 and its stereoselective directed hydrogenation of rapeseed oil. Further modification of 128 by introduction of L-type ligands yields 129, both of which act as H/D exchange catalysts.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

405

Fig. 77 Palladium and nickel Z3-allyl congeners developed by Tamm et al.

Palladium and nickel Z3-allyl complexes have also been synthesized containing these ligand sets. Pd Z3–allyl complexes were prepared via the reaction of Pd(allyl)Cl dimer with toluene solvates of 123.253 The nature of the product from these reactions depended on the solvent. When performed in THF, a bimetallic species was formed that contains a LiCl moiety bound to the metal, 130 (Fig. 77). When performed in toluene, the (NHC)PdII allyl complexes, 131 (Fig. 77), were formed. Complex 130 could be converted to 131 by dissolving in toluene, which caused the precipitation of LiCl, giving a two-step protocol. Compounds 131 were investigated as catalysts for the Buchwald–Hartwig amination of aryl halides, among which 131 (R]H; R’ ¼ H) was found to be more active. This complex was capable of high conversions to amines, but was less effective for Suzuki reactions.253 Nickel Z3–allyl complexes were synthesized using the two-step protocol via generation of NidLi bimetallic species 132 in THF, and also directly via precipitation of LiCl in toluene.254 The two-step protocol was deemed necessary as nickel Z3–allyl complexes, 133 (Fig. 77) decomposed at near room temperature via reductive elimination of zwitterionic allyl imidazolium compounds and could only be observed spectroscopically. Complex 133 could only be stabilized and crystallized in the presence of Lewis bases like CO and PR3. Attempts at using 133, and its CO and PR3 derivatives, as Buchwald–Hartwig catalysts were unsuccessful.254 Polymerization cocatalysts were also developed via incorporation of the NHC 125 onto V(V) and Ti(IV) metal centers. Reaction of 125 with Cl3V]NR led to loss of LiCl and formation of (NHC)V(]NR)Cl2, 134 (Fig. 78).255 The imido substituent was varied in complexes 134 in contrast to their previous work where different NHCs were compared. When activated by trialkylaluminum reagents, complexes 134 become highly active ethylene polymerization catalyst, especially when R ¼ o-Xyl (Fig. 78).255 Ethylene polymerization activity was probed with different aluminum cocatalysts, among which the most active was found to be triisobutylaluminum. The aluminum cocatalyst effects in ethylene–norbornene copolymerization were also investigated.256 Analogs of 134 bearing NHC 123 were also synthesized, and their polymerization activities were compared to complexes 134.257 Titanium polymerization cocatalysts were synthesized by reacting 125 with CpTiCl3 to afford CpTi(NHC)Cl2, 135 (X ¼ Cl, Fig. 78).258 When CpTiCl3 was reacted with 123, the desired analog of 135 was not formed, and instead, a low-yield product was isolated with a carbene inserting into a Cp ring. Treatment with two equivalents of MeMgBr provided CpTi(NHC)Me2, 135 (X ¼ Me, Fig. 78). Both titanium analogs were found to be competent for ethylene polymerization cocatalysts in the presence of trialkylaluminum, [Ph3C][B(C6F5)4]/trialkylaluminum, or MAO. Further, 135 was investigated as a catalyst for copolymerization of ethylene with 1-hexene.258 Five-, six-, and seven-membered cyclic NHC ligands featuring saturated backbones have been introduced by Aldridge et al.259,260 The syntheses of these complexes are generally lower yielding than the nucleophilic reactions popularized by Robinson et al.250 To our knowledge, these compounds have only been appended onto Au(I) fragments that did not demonstrate any interesting utility. Stalke and Roesky et al. demonstrated that BH3 moieties attached to the carbene carbon, 136 (R ¼ H) installed via reaction of BH3-THF with dipp carbene were robust enough to endure further deprotonation (Fig. 79).261 Reacting 136 (R ¼ H) with n-BuLi yielded the abnormal carbene–lithium adduct 137 (R ¼ H). No transition metal complexes have been formed utilizing BH3 analogs of 137.261 However, Stephan et al. developed the synthesis of the BF3 congener of 137, which may be transmetallated to give a bimetallic anionic Ag(I) complex with a lithium counterion.262 This silver complex was further transmetallated with [Ru(p-cymene) Cl]2 to afford [Ru(p-cymene)Cl(NHC)]2, 138. Reactivity of 138 was probed via salt metathesis and ligand association reactions, but these ligands have been largely unexplored beyond these Ag and Ru complexes.262

Fig. 78 Vanadium and titanium ethylene polymerization cocatalysts 134 and 135.

406

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

Fig. 79 Borate incorporation into zwitterionic imidazoliums 136, formation of abonormal adducts to lithium 137 and transmetalation procedures to afford ruthenium dimers 138.

1.12.3.3.1.1 Monodentate NHC ligands featuring other group 13 anions Beyond the borate anions discussed above, other distal anions have been appended to NHC frameworks. The initial report detailing the synthesis of the abnormal carbene highlighted the formation of a CdAl bond by the reaction of a dicarbene with AlMe3. Using the same methodology, other anionic moieties have been appended to the NHC backbone, such as Ga(CH2SiMe3)3, GaMe3 and InMe3. However, to the best of our knowledge, these ligands have never been incorporated into transition metal scaffolds.

1.12.3.3.1.2 Monodentate NHC’s featuring proximal borates Proximal borate moieties ligated to NHC frameworks were first introduced by Siebert. Zwitterionic imidazolium 139 could be synthesized by reaction of an appropriate imidazole precursor with BH3-THF. Subsequent treatment of 139 with n-BuLi yielded NHC 140, which could be treated with Mn(CO)5Br to yield Mn(I) carbonyl complex, 141 (Fig. 80).263 Homoleptic V(III) and Sc(III) complexes have been reported, as well as FeCp(CO)2 and TiCp2 fragments bearing these ligands all have been reviewed elsewhere.264 Siebert also introduced boron-, carbon-, and silicon-bridged derivatives of 139 as bidentate proximal borate platforms. When attempting to generate the biscarbene from these derivatives, it was found that the methylene-bridged imidazolium 142 was the only suitable precursor. Deprotonation of 142 resulted in the formation of carbenoid 143, the coordination chemistry of which could only be studied via trapping experiments with group IV metallocenes due to its rapid decomposition above 20  C.265 Treating 143 with Cp2MCl2 (M ¼ Ti and Zr) led to the precipitation of LiCl and formation of metallocene complex 144, which was not investigated beyond its coordination chemistry (Fig. 81). Decomposition pathways of NHC derivatives 140 and 143 typically involve the reactivity of the proximal borate moiety and are discussed in more detail here.266

Fig. 80 Siebert’s synthesis of NHC bearing proximal –BH−3 moiety, 140 and its incorporation into MnI carbonyl complex 141.

Fig. 81 Synthesis of group IV metallocenes with 143.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

1.12.4

closo-Boron and carborane anions as ligand substituents

1.12.4.1

closo-Dodecaborates as ligand substituents

407

Although not as investigated as its cousin [HCB11−], derivatives of [B12H122−] have been incorporated into ligand frameworks. This dianionic cluster has occasionally been utilized as a substituent for simple ancillary ligands. Because of its dianionic charge, its simple derivatives are often poorly soluble and hence this subfield is less developed than that of the carborane-containing ligands. When proximal to the metal’s coordination sphere, the BdH bonds in these clusters can engage in agostic-like bonding, which make them susceptible to cyclometalation, via either deprotonation or direct BdH oxidative addition.267 The term agostic-like is utilized to distinguish this mode of bonding from agostic bonding, which specifically refers to the metal bonding to CdH.

1.12.4.2

Ligands featuring proximal B12 cages

An example of a simple ligand that has been explored is the dianonic amino-closo-dodecaborate [NH2B12H2− 11] 145 cluster, which 268,269 was synthesized via electrophilic amination of [B12H2− This species has been shown to bind metals such as Au, Ru, 12] (Fig. 82). Rh, and Ni and in the cases of Rh and Ru secondary weak B-H-M interactions are observed.270 Most ligands featuring [B12H2− 12] as a ligand substituent have been discovered via mechanistic investigations into directed BdH activation and subsequent functionalization mediated by late metals. While not ancillary ligands per se, the intermediates in these transformations are ligands. For example, Duttwyler et al. have published a series of papers utilizing directed BdH activation to create novel synthons. In one report they convert the amino-closo-dodecaborate dianion 145 into the corresponding amide derivative 146, and subsequently react this compound with alkynes in the presence of a rhodium catalyst (Fig. 83). The products are vinylated cluster derivatives, with concomitant BdO bond formation to yield fused systems 147. The authors were able to isolate the s-complex intermediate 148, which features the amide ligand 145.271 It was discovered that such an alkenylation of an aryl amide closo-dodecaborate was selective for BdH activation of the ortho-position of the cluster over CdH activation of the aromatic ring with functional group tolerance on the ring that permitted subsequent cross-coupling. Even alkylation with olefins was permissible with the analogous ortho-substitution.272 A phosphine ligand containing a proximal B12 cage has also been synthesized through the cross coupling of a phosphine and a closo-dodecaborate. These form a zwitterionic diphosphine caused by BdH activation followed by a subsequent substitution of the phosphine ligand.273,274

1.12.4.3

Ligands with distal closo-dodecaborate substituents: Porphyrins, phthalocyanines, and cage complexes

Ligands that contain covalently bonded [B12H12]2− derivatives or other boron polyhedral structures are sometimes found as components of porphyrins or phthalocyanines.275–277 These are explored for Boron Neutron Capture Therapy (BNCT) as the porphyrins and phthalocyanines can accumulate in tumor cells for long periods of time, and a sufficient quantity of covalently bonded [B12H12]2− contains a high enough boron content that 10B can undergo neutron capture and subsequent decay to energetic 4 He and 7Li fission products to locally destroy a cell while leaving untargeted tissue unharmed.278,279 A series of porphyrins incorporating 1 or 2 [B12H2− 12] clusters with amino, hydroxy, and mercapto linkages have been reported. Phthalocyanine ligands carrying even more B12H2− 12 clusters were also reported, one of the designs incorporating 4 closo-dodecaborate attachments and coordinated to a zinc metal center 149 (Fig. 84).278 This was eventually extended to 8 closo-dodecaborate clusters attached to a

Fig. 82 Synthesis of amino-closo-dodecaborate 145.

Fig. 83 Double BdH activation from monofunctionalized closo-dodecaborate.

408

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

Fig. 84 Phthalocyanines incorporating 4 and 8 closo-dodecaborate structures 149 and 150, respectively.

phthalocyanine coordinated to both zinc and cobalt metal centers 150. The 8 closo-cluster complexes possessed activity from redox processes that also hinted at potential applications for electrochemical materials.280 In the same manner as porphyrins or phthalocyanines, other types of molecular scaffolds have been explored for BNCT therapy upon the installation of boron polyhedral clusters. For [B12H12]2−, one example of a cage complex has emerged with an Fe(II) clathrochelate. This molecular scaffold was modified via the nucleophilic substitution of amino-closo-dodecaborate anion.281

1.12.4.4

Multifunctionalized closo-dodecaborate in ligands

In contrast to the ligands incorporating monofunctionalized amino and hydroxyl closo-dodecaborates, examples of multifunctionalized closo-dodecaborate structures as a ligand without B-H-M interactions are extremely scarce. There is one example involving partially functionalized [B12H12-xOHx]2− (x ¼ 1–3) with titanocene dichloride to form TidO linkages from each hydroxy-containing vertex.282

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

1.12.4.5

409

closo-Carborane anions as ligand substituents

As implied in the previous section, there is far more literature related to the coordination chemistry of ligands containing closo-carborane anion substituents, perhaps because of their better solubility. The 10- and 12-vertex closo-carborane anions, 1− 2− [HCB9H1− 9 ] and [HCB11H11], are inherently more weakly coordinating than [B12H12] due to their monoanionic charge. Until recently there were very few reports of ligands featuring covalently linked closo-carborane anions.

1.12.4.6

(Alkynyl/acetylide ligands)

The first examples of the utilization of the 12-vertex carborane anions in ligand design are from Finze and coworkers.283 They reported interesting coordination chemistry with alkynyl derivatives, which were metalated at the C-terminus of the alkyne (acetylide), the p-bond of the alkyne, or both coordination modes simultaneously.283 The acetylide coordination mode can be considered as an example of a distal carborane anion ligand substituent and the p -complex binding mode can be considered an example of a ligand with a proximal closo-carborane anion group. A separate report of ferrocenyl alkyne derivatives has also been reported284 but is not described in detail here since the ferrocene is merely a spectator substituent. All of these species have been prepared via Pd catalyzed cross coupling to functionalize the C-vertex of the cluster. Finze and coworkers first reported a series of such ligands bound to Au(I) in the context of self-assembly supramolecular chemistry.283 In this work, the B-ethynyl carborane reacted with two equivalents of the corresponding R3PAuCl complex and added base to form three distinct species (Fig. 85). In these reactions, the carboranyl acetylide was formed from 153 and base, then subsequent transmetallation to R3PAuCl, and salt metathesis with another equivalent of R3PAuCl, to generate the interesting ion pair 152. In 152, the Au counteranion acts as a ligand, p-complexing the [PR3Au]+ countercation. When the phosphine ligand is large, the species remains a monomer; however, in the absence of sufficient steric congestion complexes, 152 can dimerize to form tetrametallic clusters 153, which are held together by aurophilic interactions (Fig. 85). Subsequently, Finze and coworkers found that the same ethynyl carborane salt 151 could be utilized to make a different self-assembled complexes with Ag(I).285,286 Reaction of the carboranyl alkyne with AgNO3 formed an ill-defined polymeric Ag species 154 that was subsequently utilized to form a variety of complicated complex ion pairs, as exemplified by the formation of structure 155 containing eight Ag ions (Fig. 86). Interestingly these compounds display exceptional phosphorescent properties for Ag(I) clusters. More recently, Finze and coworkers have reported that functionalizing the terminus of the alkyne functionality with an aryl or bulky silyl group allowed for the formation of simpler monomeric and dimeric Ag complexes.

1.12.4.7

closo-Carborane anions as ligand substituents for phosphines

The first phosphine containing a closo-carborane anion substituent 156 was reported in 1993 by Reed and synthesized via deprotonation of the carborane CdH vertex followed by reaction with ClPPh2 (Fig. 87). Subsequently, Reed reported the synthesis of the hexabrominated analog 157, which was prepared in a similar fashion. Subsequently, utilizing the same basic reaction

Fig. 85 The coordination chemistry of carboranyl acetylide anions.

Fig. 86 The coordination chemistry of carboranyl alkyne anions.

410

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

Fig. 87 Synthesis of anionic carboranyl phosphines.

methodology, Finze and coworkers reported a variety of other differentially functionalized anionic carboranyl phosphines, including 158. However, neither Reed nor Finze investigated these as ligands in transition metal chemistry. The first report of the implementation of such anionic phosphines in the context of transition metal chemistry and catalysis was reported by Lavallo et al. in 2013. Subsequently, Lavallo reported the first synthesis and implementation in transition metal chemistry of phosphine 159, containing the smaller 10-vertex carborane anion ligand substituent. These ligands are prepared in an analogous fashion to their 12-vertex cousins (Fig. 87). While the inductive effects of the neutral closo-carboranes H2C2B10H10/HC2B8H10 were independently investigated many years ago by Holm287 and Zhakharkin et al., the isoelectronic anionic 12- and 10-vertex carboranes’ inductive effects were a mystery until recently. Both the neutral icosahedral clusters are strong electron-withdrawing groups when C-functionalized, more so than a benzene ring. However, when a phosphine is bound to one of the B-vertices, Spokoyny et al. showed that H2C2B10H10 clusters become potent electron-donating groups. Subsequently, Lavallo and coworkers reported the first study on the inductive effects of phosphines appended with C-functionalized 12- and 10-vertex closo-carborane anions. They reported a series of anionic Rh carbonyl complexes 160 and 161 that are isoelectronic with [L2Rh(CO)2]+ compounds. Both CO complexes were prepared by the reactions of [(solv)2Rh(CO)2]+ cations generated in situ, with ligands 158 and 159 (Fig. 88). Infrared stretching frequency analysis of the compounds with 12- and 10-vertex closo-carborane anion substituents showed that they are potent electron-donating groups. The 10-vertex cluster is a stronger donor than the 12-vertex cluster due to its higher charge density, resulting from its smaller size. While the 12-vertex carborane anion substituent is similar to an isopropyl group in its electron donating ability, the 10-vertex is significantly more potent (Fig. 89). In addition to the previous study, the same group has investigated the coordinative ability of the pendant ligand substituent 162 in ligand 158. Reaction of 158 with (ClIrCOD)2 in benzene yields the unusual zwitterionic species 164 (Fig. 90).288 In the solid state, two cage BdH units interact with the metal center in an agostic-like manner. Weller289–294 and Spencer295 et al. have reported related s-complexes of the unfunctionalized anion 162 with Rh and Pt, respectively. The double BdH binding mode in 164

Fig. 88 Investigation into the inductive effects of 10 and 12-vertex carborane anions.

Fig. 89 Increasing donor ability from left to right of the o-carborane, Ph, 162, iPr, and 163.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

411

Fig. 90 Synthesis of zwitterion 164 featuring double BdH agostic “like” interactions.

conveys an extremely weak trans influence to the corresponding olefinic double bond, as the “double bond” is extremely elongated (1.451(4) A˚ ) to the point where a metallacyclopropane formalism describes the system better. This observation can be rationalized as the cluster being such a weak donor substituent that it is essentially not at all antibonding with respect to the trans-olefin, allowing better overlap for backbonding. Additionally, solution-based VT NMR studies showed that the carborane-P bond freely rotates at room temperature and has a rotational barrier of approximately 8 kcal/mol at −80 C. As mentioned in the introduction to this review, neutral carborane-based ligands have not displayed outstanding performance in catalysis. Evidence for this is that the neutral clusters are well-known to undergo cage degradation and facile BdH cyclometalation reactions. In another study, two isoelectronic Ir(I) compounds, featuring a charged carborane substituent and a neutral carborane substituent, were targeted in an effort to investigate their differing stabilities toward BdH cyclometalation. First, the zwitterionic compound 164 was converted to its anionic chloride derivative by the addition of NMe4Cl, to afford the stable Ir(I) derivative 165 (Fig. 91, left). For comparison, the isoelectronic derivative featuring a neutral carboranyl phosphine 166 was targeted (Fig. 91, right). However, the neutral isoelectronic analog to 166 could not be observed; instead, a cage BdH bond instantly underwent oxidative addition to afford the octahedral Ir(III) boryl hydride 167. This study convincingly shows that it is more difficult to BdH activate cluster 162 compared to its neutral analog. That being said, impressive catalytic BdH functionalization methodology of 162 has been reported by Weller, Duttwyler, and others.

1.12.4.8

Phosphine ligands in catalysis

In 2013, Lavallo et al. reported the first utilization of a ligand containing a closo-carborane substituent in catalysis.296 Specifically, the reported synthesis of ligand 168 and the reaction with ClAuTHT formed the zwitterionic complex 169 (Fig. 92). Notably, ligand 168 has a Tolman cone angle of 204o, which shows that it is much bulkier than P(tBu)3 (cone angle: 182 ). Nearly all previous examples of Au catalysts require an acid or Ag+ halide abstraction agent to generate a highly electrophilic Au+ catalyst. These activators add complexity to the reaction mixture and at times can act as catalysts themselves. In contrast, 169 is a single component system, which does not need an activator. Catalyst 169 displays exceptional activity for certain substrates in the hydroamination of alkynes with amines, reaching activity several orders of magnitude more active than nearly all systems for any gold-catalyzed

Fig. 91 Evidence for the cyclometalation resistance of cluster 162.

Fig. 92 Synthesis of the highly active gold catalyst 169.

412

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

reaction. It was hypothesized that this unprecedented activity is the result of both the single component nature of the catalyst, which insures rapid initiation of all catalyst molecules, and electrostatic stabilization of the Au+ center by the tethered perchlorinated carborane substituent. Subsequently it was shown that it was not possible to prepare the analogous monophosphine hydrido variant as the less sterically hindered ligand 158 affords the double phosphine-substituted version, which results in poor catalytic performance. Subsequently, the same group reported a zwitterionic Pd allyl complex 170 (Fig. 93).297 The solid-state structure revealed that a chloride of the carborane anion occupied one of the coordination sites around the desired square-planar palladium center. However, as revealed by 11B NMR, the local C5v symmetry of the carborane cluster is preserved, even when cooled to −90  C, indicating free rotation of the PdCcluster bond. This system is reminiscent of Drent-type olefin polymerization catalysts supported by anionic phosphine sulfonate ligands and thus its reactivity with simple olefins was investigated. Unfortunately, 170 does not react with ethylene (1 atm, 80 C, CD2Cl2). Furthermore, 170 rapidly isomerizes 1-hexene to its internal isomer and dimerizes. styrene at 50  C. When reacted with norbornene at 80 C, a 75% yield of polynorbornene was achieved with a PDI of 1.75 and a weight-averaged molecular weight (Mw) of 44 kDa. While the catalytic performance of 170 is rather underwhelming, it serves as proof of principle that ligands like 168 have the possibility to be utilized in olefin polymerization catalysis. The thermal stability is also noteworthy as nearly all Pd olefin polymerization catalysts rapidly decompose at 70 C or below whereas 170 does not. Later, the same authors reported a very unusual dianionic two-coordinate Pd(0) complex 171 (Fig. 94), which was accessed via reaction of ligand 158 with (TMSCH2)2PdCOD.298 When dissolved in neat Cl-C6H5, complex 171 undergoes an exceptionally fast oxidative addition reaction at ambient temperature (9 min for quantitative oxidative addition), via the intermediate 172, to yield 173 and 174 in a 9:1 ratio, respectively. The reaction was investigated both experimentally and computationally. The process begins by an essentially barrierless ligand dissociation of one phosphine ligand 158 and subsequent coordination of the arene. Then a low energy barrier, oxidative addition (4 kcal/mol) produces 172. The intermediate 172 contains an aryl ligand as well as a phosphine 158, which is stabilized by agostic-like interactions with the.

Fig. 93 Synthesis of the zwitterionic Pd complex 170.

Fig. 94 Rapid Oxidative Addition of Cl-C6H5.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

413

BdH cage. The desirable monoanionic compound 173 formed via reassociation of 158 and simultaneously eliminated LiCl. Compound 174 arises from 172 via a divergent s-bond metathesis of a BdH bond with the aryl ligand to produce benzene, followed by reassociation of 158. Comparatively, Pd(PCy3)2 afforded 70% conversion in 24 h and Pd(P(tBu)3)2 showed no reactivity at ambient temperature.298 The remarkable reactivity of 171 is thought to occur because two negatively charged phosphines bound to the zerovalent Pd center creates a situation where coulombic repulsion renders ligand dissociation very favorable. The result is rapid access to a monophosphine-ligated Pd(0) intermediate, which readily undergoes oxidative addition at ambient temperature. To highlight the difference in electronic properties of CB11 substituents and isoelectronic C2B10 substituents, compound 175 was prepared (Fig. 95). This neutral dicoordinate Pd(0) species does not undergo oxidative addition with Cl-C6H5, even at elevated temperatures. Further, the authors found compound 171 to be a competent catalyst for the Kumada coupling of simple chloroarenes, but the activity of the catalyst was nowhere near the best catalysts for the same reaction. The lower activity, relative to the fast oxidative addition, was likely the result of slow transmetallation/reductive elimination steps or simply catalyst deactivation via similar s-bond metathesis pathways that form 174. The utilization of BdH functionalized cluster substituents that prevent cyclometalation may result in more active catalysts. Subsequently, Lavallo and Jordan reported the synthesis and implementation of perchlorinated phosphines 176–178 in the exploration of Pd polymerization catalysts (Fig. 96).299 Reaction of phosphines 176 and 178 with ClPd(Me)COD and subsequent reaction with AgBF4 led to the formation of the zwitterionic methyl complexes 179 and 180, respectively. In the solid state, a chloride from the cluster makes a weak bonding interaction with the Pd-center. In contrast, when ligand 178 was employed, complex 181 was formed, which features a o-methoxy chelate that prevented interaction with the carborane cage Cl substituents. Compounds showed moderate to good activity for the oligomerization of ethylene while 181 is an ethylene polymerization catalyst. Most recently it was shown that when compound 180 was heated or photolyzed, it eliminated ethane producing the very unusual Pd(I) dimer 182. Complex 182 is unusual as it is nearly an unsupported L-Pd-Pd-L fragment as both the p-arene and PddCl interactions are weak. When a THF solution of 182 was exposed to ethylene, a precipitate immediately formed. Upon dissolving the precipitate in solvents where it is soluble, such as DCM, ethylene is released and reformed 182, making it impossible to characterize the product by solution-state NMR. Through a combination of solid-state NMR and micro-electron diffraction (micro-ED) techniques, the structure was determined to be that of the 1,2-dipalladated Pd(II) species 183 (Fig. 97).300 This is the first unambiguous determination of a 1,2-dipalladation reaction and also the first utilization of micro-ED to identify an unknown organometallic species.

Fig. 95 Compound 175, which is isoelectronic and steric with 171 does not react with ClC6H5.

Fig. 96 Synthesis of zwitterionic Pd methyl complexes supported by phosphines containing perchlorinated 10-vertex carborane anions.

414

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

Fig. 97 Synthesis of a “naked” Pd(I) dimer and its reversible reaction with ethylene.

1.12.4.9

N-Heterocyclic carbene ligands with anionic carborane substituents

Recently, Lavallo et al. developed synthetic methodology to access several classes of NHCs that contain carborane clusters directly bound to the nitrogen atom of the NHC ring.301 This section highlights the synthesis and unique behavior of these species, as well as the preliminary coordination chemistry and catalysis of these ligands with Au(I). This started with the preparation of the anionic imidazolium salt 184, which is the conjugate acid of the targeted NHCs. Depending on the conditions and base employed, the imidazolium anion 184 is a suitable precursor to three unique NHC lithium salts 185–186 (Fig. 98). Selective C-2 deprotonation of 184 could be achieved using lithium bis(trimethylsilyl)amide (LiHMDS) to afford the so called “normal” carbene 185. Selective C-5 deprotonation to form the “abnormal” NHC 186 was accomplished by using lithium diisopropylamide (LDA) at −78  C. NHCs 185 and 186 are the first examples of two NHC constitutional isomers derived from a single precursor. The abnormal C-5 isomer 186 is higher in energy than 185 and slowly isomerizes to the normal C-2 species over time. Imidazolium anion 184 could also be doubly deprotonated with n-BuLi to form the trianionic C-2/C-5 lithium salt 187. The selective formation of three different NHCs from a single precursor is not possible with standard hydrocarbon substituents and is likely the result of both electronic and steric effects conveyed by the cluster groups. In a separate report, the same group demonstrated the synthesis of unsymmetrical NHCs featuring only one carborane substituent (Fig. 99).302 When the zwitterionic imidazolium compound 188 was treated with LiHMDS, the monoanionic Li carbenoid 177 was produced. When 188 was reacted with 2 equivalents of n-BuLi, double deprotonation was achieved, forming

Fig. 98 Synthesis of different polyanionic carbenoids appended with 12-vertex cluster anions.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

415

Fig. 99 Synthesis of unsymmetrical Li-carbenoids.

Fig. 100 Synthesis of Au(I) complexes 191 and 192.

the dianionic Li complex 190. However, in contrast to the anionic imidazolium salt 184 featuring two carborane substituents, it was not possible to selectively prepare the abnormal C-5 isomer from 188. Subsequently, it was demonstrated that carbenoids 185 and 189 could be transmetallated to Au(I) to produce the anionic and zwitterionic complexes 179 and 180, respectively (Fig. 100).303 Structural studies of these complexes revealed that the 12-vertex carborane ligand substituent conveys a 3.7% greater percent buried volume compared to an adamantyl group. Most recently, Lavallo and coworkers introduced a new way to utilize such species, as strongly coordinating ligands to form weakly coordinating yet functional anionic organometallic compounds.304 By binding the dianionic carbenoids to Au(I), one can view the entire -ate complex as a weakly coordinating anion. Indeed, they demonstrated this by first reacting 191 with AgBF4 to form the new heterobimetallic Au−/Ag+ ion pair 193 (Fig. 101). After decoordination of the THF solvent with acetonitrile, they reacted 193 with trityl chloride to form the corresponding trityl salt 194. Compound 194 was only stable in acetonitrile, which limits its utility as a hydride or alkide abstraction agent. In other solvents, it undergoes both cyclometallation reactions at a BdH vertex and electrophilic arylation of the cage with the trityl cation. The cyclometalated compound is notable because it is the only example of a Au(III) compound with two cis-accessible coordination sites. To block such cyclometallation and electrophilic arylation decomposition pathways, they devised the synthetic methodology to create NHCs with polyhalogenated carborane substituents. These are indeed dianionic NHCs and not carbenoids as shown by NMR and single crystal X-ray diffraction studies. Following similar protocol as described in Fig. 101, the polyhalogenated ion pairs 195 and 196 can be prepared (Fig. 102). Indeed, the trityl salts 196 did not undergo any cyclometalation reactions or electrophilic arylation making them much more useful reagents. In the same report, it was shown that 195 and 196 could be utilized as reagents for Ag salt metathesis and hydride abstractions to form the new heterobimetallic Ag−/Ir+ ion pairs 197 and silylium salts 198 (Fig. 103).304 The silylium salts should be useful reagents for halide abstractions from inorganic or organic compounds to make an array of exotic salts. After clearly demonstrating that these organometallic species are weakly coordinating anions that can be paired with very reactive countercations, the authors demonstrated that these gold anions are not simply spectators. They showed this by utilizing the Au anions in hydroamination catalysis of

Fig. 101 Synthesis of heterobimetallic transition metal and hybrid transition metal main-group ion pairs.

416

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

Fig. 102 Stable Ag+ and trityl salts of functional weakly coordinating organometallic anions.

Fig. 103 Heterobimetallic Au−/Ir+ and hybrid Au−/SiEt+3 .

alkynes and amines, where they exhibited turnover numbers approaching 100,000 with some substrates. Thus, this new class of molecules are functional weakly coordinating anions. This finding has implications in tandem/cooperative catalysis and small molecule activation.

1.12.5

Concluding remarks

This chapter outlines the rich chemistry of ligands containing moderate to weakly coordinating anions. Whether of fundamental interest, catalysis, or materials chemistry, it is clear that this is a rewarding approach to ligand design. The future looks bright for this diverse array of chemistry, and new surprises keep emerging as exemplified by the emergence of functional weakly coordinating anions described in the end of the chapter.

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.

Strauss, S. H. Chem. Rev. 1993, 93, 927–942. Engesser, T. A.; Lichtenthaler, M. R.; Schleep, M.; Krossing, I. Chem. Soc. Rev. 2016, 45, 789–899. Reed, C. A. Acc. Chem. Res. 2010, 43, 121–128. Fisher, S. P.; Tomich, A. W.; Lovera, S. O.; Kleinsasser, J. F.; Guo, J.; Asay, M. J.; Nelson, H. M.; Lavallo, V. Chem. Rev. 2019, 119, 8262–8290. Olah, G. A., Prakash, G. K. S., et al. In Carbocation Chemistry; John Wiley & Sons, Inc., 2004 Schrock, R. R.; Osborn, J. A. J. Am. Chem. Soc. 1976, 98, 2143–2147. Riddlestone, I. M.; Kraft, A.; Schaefer, J.; Krossing, I. Angew. Chem. Int. Ed. 2018, 57, 13982–14024. Nishida, H.; Takada, N.; Yoshimura, M.; Sonoda, T.; Kobayashi, H. Bull. Chem. Soc. Jpn. 1984, 57, 2600–2604. Massey, A. G.; Park, A. J.; Stone, F. G. A. Proc. Chem. Soc., London 1963, 212. Axtell, J. C.; Saleh, L. M. A.; Qian, E. A.; Wixtrom, A. I.; Spokoyny, A. M. Inorg. Chem. 2018, 57, 2333–2350. Jay, R.; Tomich, A. W.; Zhang, J.; Zhao, Y.; De Gorostiza, A.; Lavallo, V.; Guo, J. ACS Appl. Mater. Interfaces 2019, 11, 11414–11420. Pitochelli, A. R.; Hawthorne, F. M. J. Am. Chem. Soc. 1960, 82, 3228–3229. Knoth, W. H. J. Am. Chem. Soc. 1967, 89, 1274–1275. Douvris, C.; Michl, J. Chem. Rev. 2013, 113, PR179–PR233. Farha, O. K.; Julius, R. L.; Lee, M. W.; Huertas, R. E.; Knobler, C. B.; Hawthorne, M. F. J. Am. Chem. Soc. 2005, 127, 18243–18251. Drent, E.; van Dijk, R.; van Ginkel, R.; van Oort, B.; Pugh, R. I. Chem. Commun. 2002. https://doi.org/10.1039/B111629K, 964-965. Hearley, A. K.; Nowack, R. J.; Rieger, B. Organometallics 2005, 24, 2755–2763. Zhang, D.; Guironnet, D.; Göttker-Schnetmann, I.; Mecking, S. Organometallics 2009, 28, 4072–4078. Kochi, T.; Yoshimura, K.; Nozaki, K. Dalton Trans. 2006. https://doi.org/10.1039/B512452M, 25-27. Kochi, T.; Noda, S.; Yoshimura, K.; Nozaki, K. J. Am. Chem. Soc. 2007, 129, 8948–8949. Ito, S.; Munakata, K.; Nakamura, A.; Nozaki, K. J. Am. Chem. Soc. 2009, 131, 14606–14607. Ito, S.; Kanazawa, M.; Munakata, K.; Kuroda, J.-I.; Okumura, Y.; Nozaki, K. J. Am. Chem. Soc. 2011, 133, 1232–1235. Luo, S.; Vela, J.; Lief, G. R.; Jordan, R. F. J. Am. Chem. Soc. 2007, 129, 8946–8947. Weng, W.; Shen, Z.; Jordan, R. F. J. Am. Chem. Soc. 2007, 129, 15450–15451. Shen, Z.; Jordan, R. F. Macromolecules 2010, 43, 8706–8708. Cai, Z.; Shen, Z.; Zhou, X.; Jordan, R. F. ACS Catalysis 2012, 2, 1187–1195. Wucher, P.; Goldbach, V.; Mecking, S. Organometallics 2013, 32, 4516–4522.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions 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. 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.

417

Schuster, N.; Rünzi, T.; Mecking, S. Macromolecules 2016, 49, 1172–1179. Guironnet, D.; Caporaso, L.; Neuwald, B.; Göttker-Schnetmann, I.; Cavallo, L.; Mecking, S. J. Am. Chem. Soc. 2010, 132, 4418–4426. Jian, Z.; Wucher, P.; Mecking, S. Organometallics 2014, 33, 2879–2888. Guironnet, D.; Roesle, P.; Rünzi, T.; Göttker-Schnetmann, I.; Mecking, S. J. Am. Chem. Soc. 2009, 131, 422–423. Friedberger, T.; Wucher, P.; Mecking, S. J. Am. Chem. Soc. 2012, 134, 1010–1018. Bouilhac, C.; Rünzi, T.; Mecking, S. Macromolecules 2010, 43, 3589–3590. Berkefeld, A.; Mecking, S. Angew. Chem. Int. Ed. 2008, 47, 2538–2542. Leicht, H.; Göttker-Schnetmann, I.; Mecking, S. Angew. Chem. Int. Ed. 2013, 52, 3963–3966. Lanzinger, D.; Giuman, M. M.; Anselment, T. M. J.; Rieger, B. ACS Macro Lett. 2014, 3, 931–934. Nowack, R. J.; Hearley, A. K.; Rieger, B. Z. Anorg. Allg. Chem. 2005, 631, 2775–2781. Zhou, X.; Bontemps, S.; Jordan, R. F. Organometallics 2008, 27, 4821–4824. Noda, S.; Kochi, T.; Nozaki, K. Organometallics 2009, 28, 656–658. Perrotin, P.; McCahill, J. S. J.; Wu, G.; Scott, S. L. Chem. Commun. 2011, 47, 6948–6950. Chen, M.; Zou, W.; Cai, Z.; Chen, C. Polym. Chem. 2015, 6, 2669–2676. Wu, Z.; Chen, M.; Chen, C. Organometallics 2016, 35, 1472–1479. Zhang, D.; Chen, C. Angew. Chem. Int. Ed. 2017, 56, 14672–14676. Chen, M.; Chen, C. ACS Catalysis 2017, 7, 1308–1312. Chen, M.; Yang, B.; Chen, C. Angew. Chem. Int. Ed. 2015, 54, 15520–15524. Zábranský, M.; Císarˇová, I.; Trzeciak, A. M.; Alsalahi, W.; Štepnicka, P. Organometallics 2019, 38, 479–488. Szabo, M. J.; Galea, N. M.; Michalak, A.; Yang, S.-Y.; Groux, L. F.; Piers, W. E.; Ziegler, T. J. Am. Chem. Soc. 2005, 127, 14692–14703. Szabo, M. J.; Galea, N. M.; Michalak, A.; Yang, S.-Y.; Groux, L. F.; Piers, W. E.; Ziegler, T. Organometallics 2005, 24, 2147–2156. Szabo, M. J.; Jordan, R. F.; Michalak, A.; Piers, W. E.; Weiss, T.; Yang, S.-Y.; Ziegler, T. Organometallics 2004, 23, 5565–5572. Wang, L.; Carrow, B. P. ACS Catalysis 2019, 9, 6821–6836. Hopkinson, M. N.; Richter, C.; Schedler, M.; Glorius, F. Nature 2014, 510, 485–496. Nagai, Y.; Kochi, T.; Nozaki, K. Organometallics 2009, 28, 6131–6134. Li, M.; Song, H.; Wang, B. Eur. J. Inorg. Chem. 2015, 2015, 4055–4061. Brown, M. K.; May, T. L.; Baxter, C. A.; Hoveyda, A. H. Angew. Chem. Int. Ed. 2007, 46, 1097–1100. Hoveyda, A. H.; Malcolmson, S. J.; Meek, S. J.; Zhugralin, A. R. Angew. Chem. Int. Ed. 2010, 49, 34–44. Gillingham, D. G.; Hoveyda, A. H. Angew. Chem. Int. Ed. 2007, 46, 3860–3864. Zhou, X.; Jordan, R. F. Organometallics 2011, 30, 4632–4642. Trofimenko, S. J. Am. Chem. Soc. 1966, 88, 1842–1844. Trofimenko, S. Chem. Rev. 1972, 72, 497–509. Trofimenko, S. Acc. Chem. Res. 1971, 4, 17–22. Trofimenko, S. Chem. Rev. 1993, 93, 943–980. Muñoz-Molina, J. M.; Belderrain, T. R.; Pérez, P. J. Coord. Chem. Rev. 2019, 390, 171–189. Riordan, C. G. Coord. Chem. Rev. 2010, 254, 1815–1825. Thomas, J. C.; Peters, J. C. Inorg. Chem. 2003, 42, 5055–5073. Betley, T. A.; Peters, J. C. Inorg. Chem. 2002, 41, 6541–6543. Kernbach, U.; Ramm, M.; Luger, P.; Fehlhammer, W. P. Angew. Chem. Int. Ed. Engl. 1996, 35, 310–312. Dias, H. V. R.; Wang, X.; Diyabalanage, H. V. K. Inorg. Chem. 2005, 44, 7322–7324. Hsu, S. C. N.; Chang, Y.-L.; Chuang, W.-J.; Chen, H.-Y.; Lin, I. J.; Chiang, M. Y.; Kao, C.-L.; Chen, H.-Y. Inorg. Chem. 2012, 51, 9297–9308. Díaz-Requejo, M. M.; Belderraín, T. R.; Nicasio, M. C.; Trofimenko, S.; Pérez, P. J. J. Am. Chem. Soc. 2003, 125, 12078–12079. Santini, C.; Pellei, M.; Lobbia, G. G.; Papini, G. Mini Rev. Org. Chem. 2010, 7, 84–124. Pellei, M.; Gioia Lobbia, G.; Papini, G.; Santini, C. Mini Rev. Org. Chem. 2010, 7, 173–203. Chia, L. M. L.; Radojevic, S.; Scowen, I. J.; McPartlin, M.; Halcrow, M. A. Dalton Trans. 2000. https://doi.org/10.1039/A907258F, 133-140. Gutierrez, E.; Monge, A.; Nicasio, M. C.; Poveda, M. L.; Carmona, E. J. Am. Chem. Soc. 1994, 116, 791–792. Lee, H.; Jordan, R. F. J. Am. Chem. Soc. 2005, 127, 9384–9385. Jayaratna, N. B.; Pardue, D. B.; Ray, S.; Yousufuddin, M.; Thakur, K. G.; Cundari, T. R.; Dias, H. V. R. Dalton Trans. 2013, 42, 15399–15410. Thomas, J. C.; Peters, J. C. J. Am. Chem. Soc. 2003, 125, 8870–8888. Foley, N. A.; Lee, J. P.; Ke, Z.; Gunnoe, T. B.; Cundari, T. R. Acc. Chem. Res. 2009, 42, 585–597. Meucci, E. A.; Nguyen, S. N.; Camasso, N. M.; Chong, E.; Ariafard, A.; Canty, A. J.; Sanford, M. S. J. Am. Chem. Soc. 2019, 141, 12872–12879. Bergman, R. G. Science 1984, 223, 902–908. Jones, W. D.; Hessell, E. T. J. Am. Chem. Soc. 1992, 114, 6087–6095. Wick, D. D.; Reynolds, K. A.; Jones, W. D. J. Am. Chem. Soc. 1999, 121, 3974–3983. Northcutt, T. O.; Wick, D. D.; Vetter, A. J.; Jones, W. D. J. Am. Chem. Soc. 2001, 123, 7257–7270. Vetter, A. J.; Flaschenriem, C.; Jones, W. D. J. Am. Chem. Soc. 2005, 127, 12315–12322. Evans, M. E.; Burke, C. L.; Yaibuathes, S.; Clot, E.; Eisenstein, O.; Jones, W. D. J. Am. Chem. Soc. 2009, 131, 13464–13473. Evans, M. E.; Li, T.; Jones, W. D. J. Am. Chem. Soc. 2010, 132, 16278–16284. Jiao, Y.; Morris, J.; Brennessel, W. W.; Jones, W. D. J. Am. Chem. Soc. 2013, 135, 16198–16212. Jiao, Y.; Evans, M. E.; Morris, J.; Brennessel, W. W.; Jones, W. D. J. Am. Chem. Soc. 2013, 135, 6994–7004. Procacci, B.; Jiao, Y.; Evans, M. E.; Jones, W. D.; Perutz, R. N.; Whitwood, A. C. J. Am. Chem. Soc. 2015, 137, 1258–1272. Guan, J.; Wriglesworth, A.; Sun, X. Z.; Brothers, E. N.; Zaric, S. D.; Evans, M. E.; Jones, W. D.; Towrie, M.; Hall, M. B.; George, M. W. J. Am. Chem. Soc. 2018, 140, 1842–1854. Fructos, M. R.; Trofimenko, S.; Díaz-Requejo, M. M.; Pérez, P. J. J. Am. Chem. Soc. 2006, 128, 11784–11791. Shay, D. T.; Yap, G. P. A.; Zakharov, L. N.; Rheingold, A. L.; Theopold, K. H. Angew. Chem. 2005, 44, 1508–1510. Hikichi, S.; Yoshizawa, M.; Sasakura, Y.; Akita, M.; Moro-oka, Y. J. Am. Chem. Soc. 1998, 120, 10567–10568. Sallmann, M.; Limberg, C. Acc. Chem. Res. 2015, 48, 2734–2743. Kitajima, N.; Fujisawa, K.; Fujimoto, C.; Morooka, Y.; Hashimoto, S.; Kitagawa, T.; Toriumi, K.; Tatsumi, K.; Nakamura, A. J. Am. Chem. Soc. 1992, 114, 1277–1291. Mealli, C.; Arcus, C. S.; Wilkinson, J. L.; Marks, T. J.; Ibers, J. A. J. Am. Chem. Soc. 1976, 98, 711–718. Chen, C.; Lee, H.; Jordan, R. F. Organometallics 2010, 29, 5373–5381. Dias, H. V. R.; Jin, W. Inorg. Chem. 2000, 39, 815–819. Dias, H. V. R.; Thankamani, J. Acta Crystallogr. C 2013, 69, 959–962. Dias, H. V. R.; Browning, R. G.; Polach, S. A.; Diyabalanage, H. V. K.; Lovely, C. J. J. Am. Chem. Soc. 2003, 125, 9270–9271.

418

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions

100. Dias, H. V. R.; Browning, R. G.; Richey, S. A.; Lovely, C. J. Organometallics 2004, 23, 1200–1202. 101. Caballero, A.; Despagnet-Ayoub, E.; Mar Díaz-Requejo, M.; Díaz-Rodríguez, A.; González-Núñez, M. E.; Mello, R.; Muñoz, B. K.; Ojo, W.-S.; Asensio, G.; Etienne, M.; Pérez, P. J. Science 2011, 332, 835. 102. Ponduru, T. T.; Sun, Z.; Cundari, T. R.; Rasika Dias, H. V. ChemCatChem 2019, 11, 4966–4973. 103. Arenas, I.; Fuentes, M.Á.; Álvarez, E.; Díaz, Y.; Caballero, A.; Castillón, S.; Pérez, P. J. Inorg. Chem. 2014, 53, 3991–3999. 104. Pettinari, C.; Cingolani, A.; Lobbia, G. G.; Marchetti, F.; Martini, D.; Pellei, M.; Pettinari, R.; Santini, C. Polyhedron 2004, 23, 451–469. 105. Cao, H.-J.; Zhao, Q.; Zhang, Q.-F.; Li, J.; Hamilton, E. J. M.; Zhang, J.; Wang, L.-S.; Chen, X. Dalton Trans. 2016, 45, 10194–10199. 106. Muñoz-Hernández, M.-Á.; Montiel-Palma, V. Inorg. Chim. Acta 2009, 362, 4328–4339. 107. Riddlestone, I. M.; Keller, S.; Kirschenmann, F.; Schorpp, M.; Krossing, I. Eur. J. Inorg. Chem. 2019, 2019, 59–67. 108. Klahn, M.; Fischer, C.; Spannenberg, A.; Rosenthal, U.; Krossing, I. Tetrahedron Lett. 2007, 48, 8900–8903. 109. Breakell, K. R.; Patmore, D. J.; Storr, A. J. Chem. Soc. Dalton Trans. 1975. https://doi.org/10.1039/DT9750000749, 749-754. 110. Cortes-Llamas, S.; Velázquez-Carmona, M.-Á.; Muñoz-Hernández, M.-Á. Inorg. Chem. Commun. 2005, 8, 155–158. 111. Snyder, C. J.; Heeg, M. J.; Winter, C. H. Inorg. Chem. 2011, 50, 9210–9212. 112. Dunne, J. F.; Su, J.; Ellern, A.; Sadow, A. D. Organometallics 2008, 27, 2399–2401. 113. Reinig, R. R.; Mukherjee, D.; Weinstein, Z. B.; Xie, W.; Albright, T.; Baird, B.; Gray, T. S.; Ellern, A.; Miller, G. J.; Winter, A. H.; Bud’ko, S. L.; Sadow, A. D. Eur. J. Inorg. Chem. 2016, 2016, 2486–2494. 114. Fraile, J. M.; García, J. I.; Mayoral, J. A. Coord. Chem. Rev. 2008, 252, 624–646. 115. Mazet, C.; Köhler, V.; Pfaltz, A. Angew. Chem. 2005, 44, 4888–4891. 116. Narwane, M.; Chang, Y.-L.; Ching, W.-M.; Tsai, M.-L.; Hsu, S. C. N. Inorg. Chim. Acta 2019, 495, 118966. 117. Jeong, S. Y.; Lalancette, R. A.; Lin, H.; Lupinska, P.; Shipman, P. O.; John, A.; Sheridan, J. B.; Jäkle, F. Inorg. Chem. 2016, 55, 3605–3615. 118. Conley, B. L.; Guess, D.; Williams, T. J. J. Am. Chem. Soc. 2011, 133, 14212–14215. 119. Pawar, G. M.; Sheridan, J. B.; Jäkle, F. Eur. J. Inorg. Chem. 2016, 2016, 2227–2235. 120. Janiak, C. ChemComm 1994. https://doi.org/10.1039/C39940000545, 545-547. 121. Peters, J. C.; Feldman, J. D.; Tilley, T. D. J. Am. Chem. Soc. 1999, 121, 9871–9872. 122. Barney, A. A.; Heyduk, A. F.; Nocera, D. G. Chem. Commun. 1999. https://doi.org/10.1039/a906560a, 2379-2380. 123. Turculet, L.; Feldman, J. D.; Tilley, T. D. Organometallics 2004, 23, 2488–2502. 124. Feldman, J. D.; Peters, J. C.; Tilley, T. D. Organometallics 2002, 21, 4050–4064. 125. Jiménez, S.; López, J. A.; Ciriano, M. A.; Tejel, C.; Martínez, A.; Sánchez-Delgado, R. A. Organometallics 2009, 28, 3193–3202. 126. Tejel, C.; Ciriano, M. A.; Passarelli, V. Chem. A Eur. J. 2011, 17, 91–95. 127. Betley, T. A.; Peters, J. C. Angew. Chem. Int. Ed. 2003, 42, 2385–2389. 128. Jenkins, D. M.; Di Bilio, A. J.; Allen, M. J.; Betley, T. A.; Peters, J. C. J. Am. Chem. Soc. 2002, 124, 15336–15350. 129. Jenkins, D. M.; Betley, T. A.; Peters, J. C. J. Am. Chem. Soc. 2002, 124, 11238–11239. 130. Betley, T. A.; Peters, J. C. Inorg. Chem. 2003, 42, 5074–5084. 131. Christopher Thomas, J.; Peters, J. C. Polyhedron 2004, 23, 2901–2913. 132. McCormick, T.; Jia, W.-L.; Wang, S. Inorg. Chem. 2006, 45, 147–155. 133. Ángeles Fuentes, M.; Álvarez, E.; Caballero, A.; Pérez, P. J. Organometallics 2012, 31, 959–965. 134. McCain, M. N.; Schneider, S.; Salata, M. R.; Marks, T. J. Inorg. Chem. 2008, 47, 2534–2542. 135. Walker, J. M.; Cox, A. M.; Wang, R.; Spivak, G. J. Organometallics 2010, 29, 6121–6124. 136. Lipke, M. C.; Tilley, T. D. J. Am. Chem. Soc. 2011, 133, 16374–16377. 137. Lipke, M. C.; Neumeyer, F.; Tilley, T. D. J. Am. Chem. Soc. 2014, 136, 6092–6102. 138. Saouma, C. T.; Müller, P.; Peters, J. C. J. Am. Chem. Soc. 2009, 131, 10358–10359. 139. Saouma, C. T.; Lu, C. C.; Peters, J. C. Inorg. Chem. 2012, 51, 10043–10054. 140. Brown, S. D.; Peters, J. C. J. Am. Chem. Soc. 2005, 127, 1913–1923. 141. Brown, S. D.; Mehn, M. P.; Peters, J. C. J. Am. Chem. Soc. 2005, 127, 13146–13147. 142. Brown, S. D.; Betley, T. A.; Peters, J. C. J. Am. Chem. Soc. 2003, 125, 322–323. 143. Brown, S. D.; Peters, J. C. J. Am. Chem. Soc. 2004, 126, 4538–4539. 144. Betley, T. A.; Peters, J. C. J. Am. Chem. Soc. 2004, 126, 6252–6254. 145. Daida, E. J.; Peters, J. C. Inorg. Chem. 2004, 43, 7474–7485. 146. Turculet, L.; Feldman, J. D.; Tilley, T. D. Organometallics 2003, 22, 4627–4629. 147. Hendrich, M. P.; Gunderson, W.; Behan, R. K.; Green, M. T.; Mehn, M. P.; Betley, T. A.; Lu, C. C.; Peters, J. C. Proc. Natl. Acad. Sci. 2006, 103, 17107. 148. Lu, C. C.; Saouma, C. T.; Day, M. W.; Peters, J. C. J. Am. Chem. Soc. 2007, 129, 4–5. 149. MacBeth, C. E.; Thomas, J. C.; Betley, T. A.; Peters, J. C. Inorg. Chem. 2004, 43, 4645–4662. 150. Lu, C. C.; Peters, J. C. Inorg. Chem. 2006, 45, 8597–8607. 151. Handford, R. C.; Smith, P. W.; Tilley, T. D. Organometallics 2018, 37, 4077–4085. 152. Thomas, J. C.; Peters, J. C. J. Am. Chem. Soc. 2001, 123, 5100–5101. 153. Lu, C. C.; Peters, J. C. J. Am. Chem. Soc. 2004, 126, 15818–15832. 154. Lu, C. C.; Peters, J. C. J. Am. Chem. Soc. 2002, 124, 5272–5273. 155. Wu, F.; Jordan, R. F. Organometallics 2006, 25, 5631–5637. 156. Fischer, P. J.; Avena, L.; Bohrmann, T. D.; Neary, M. C.; Putka, G. K.; Sullivan, K. P. Organometallics 2014, 33, 1300–1309. 157. Fischer, P. J.; Weberg, A. B.; Bohrmann, T. D.; Xu, H.; Young, V. G. Dalton Trans. 2015, 44, 3737–3744. 158. Casado, M. A.; Hack, V.; Camerano, J. A.; Ciriano, M. A.; Tejel, C.; Oro, L. A. Inorg. Chem. 2005, 44, 9122–9124. 159. Thomas, C. M.; Mankad, N. P.; Peters, J. C. J. Am. Chem. Soc. 2006, 128, 4956–4957. 160. Sircoglou, M.; Bouhadir, G.; Saffon, N.; Miqueu, K.; Bourissou, D. Organometallics 2008, 27, 1675–1678. 161. Sircoglou, M.; Saffon, N.; Miqueu, K.; Bouhadir, G.; Bourissou, D. Organometallics 2013, 32, 6780–6784. 162. Devillard, M.; Nicolas, E.; Appelt, C.; Backs, J.; Mallet-Ladeira, S.; Bouhadir, G.; Slootweg, J. C.; Uhl, W.; Bourissou, D. Chem. Commun. 2014, 50, 14805–14808. 163. Santini, C.; Marinelli, M.; Pellei, M. Eur. J. Inorg. Chem. 2016, 2016, 2312–2331. 164. Muñoz, S. B.; Foster, W. K.; Lin, H.-J.; Margarit, C. G.; Dickie, D. A.; Smith, J. M. Inorg. Chem. 2012, 51, 12660–12668. 165. Martinez, J. L.; Lee, W.-T.; Pink, M.; Chen, C.-H.; Smith, J. M. Polyhedron 2018, 156, 297–302. 166. Fränkel, R.; Birg, C.; Kernbach, U.; Habereder, T.; Nöth, H.; Fehlhammer, W. P. Angew. Chem. Int. Ed. 2001, 40, 1907–1910. 167. Fränkel, R.; Kniczek, J.; Ponikwar, W.; Nöth, H.; Polborn, K.; Fehlhammer, W. P. Inorg. Chim. Acta 2001, 312, 23–39. 168. Hickey, A. K.; Chen, C.-H.; Pink, M.; Smith, J. M. Organometallics 2015, 34, 4560–4566. 169. Hickey, A. K.; Lee, W.-T.; Chen, C.-H.; Pink, M.; Smith, J. M. Organometallics 2016, 35, 3069–3073. 170. Hickey, A. K.; Lutz, S. A.; Chen, C.-H.; Smith, J. M. Chem. Commun. 2017, 53, 1245–1248.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions 171. 172. 173. 174. 175. 176. 177. 178. 179. 180. 181. 182. 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. 233. 234. 235. 236. 237. 238. 239. 240.

419

Smith, J. M.; Subedi, D. Dalton Trans. 2012, 41, 1423–1429. Scepaniak, J. J.; Fulton, M. D.; Bontchev, R. P.; Duesler, E. N.; Kirk, M. L.; Smith, J. M. J. Am. Chem. Soc. 2008, 130, 10515–10517. Scepaniak, J. J.; Young, J. A.; Bontchev, R. P.; Smith, J. M. Angew. Chem. Int. Ed. 2009, 48, 3158–3160. Crandell, D. W.; Muñoz, S. B.; Smith, J. M.; Baik, M.-H. Chem. Sci. 2018, 9, 8542–8552. Muñoz Iii, S. B.; Lee, W.-T.; Dickie, D. A.; Scepaniak, J. J.; Subedi, D.; Pink, M.; Johnson, M. D.; Smith, J. M. Angew. Chem. Int. Ed. 2015, 54, 10600–10603. Lee, W.-T.; Juarez, R. A.; Scepaniak, J. J.; Muñoz, S. B.; Dickie, D. A.; Wang, H.; Smith, J. M. Inorg. Chem. 2014, 53, 8425–8430. Lin, H.-J.; Siretanu, D.; Dickie, D. A.; Subedi, D.; Scepaniak, J. J.; Mitcov, D.; Clérac, R.; Smith, J. M. J. Am. Chem. Soc. 2014, 136, 13326–13332. Scepaniak, J. J.; Bontchev, R. P.; Johnson, D. L.; Smith, J. M. Angew. Chem. Int. Ed. 2011, 50, 6630–6633. Valdez-Moreira, J. A.; Millikan, S. P.; Gao, X.; Carta, V.; Chen, C.-H.; Smith, J. M. Inorg. Chem. 2020, 59, 579–583. Martinez, J. L.; Lin, H.-J.; Lee, W.-T.; Pink, M.; Chen, C.-H.; Gao, X.; Dickie, D. A.; Smith, J. M. J. Am. Chem. Soc. 2017, 139, 14037–14040. Mathonière, C.; Lin, H.-J.; Siretanu, D.; Clérac, R.; Smith, J. M. J. Am. Chem. Soc. 2013, 135, 19083–19086. Ding, M.; Rouzières, M.; Losovyj, Y.; Pink, M.; Clérac, R.; Smith, J. M. Inorg. Chem. 2015, 54, 9075–9080. Scepaniak, J. J.; Vogel, C. S.; Khusniyarov, M. M.; Heinemann, F. W.; Meyer, K.; Smith, J. M. Science 2011, 331, 1049. Cutsail Iii, G. E.; Stein, B. W.; Subedi, D.; Smith, J. M.; Kirk, M. L.; Hoffman, B. M. J. Am. Chem. Soc. 2014, 136, 12323–12336. Martinez, J. L.; Lutz, S. A.; Beagan, D. M.; Gao, X.; Pink, M.; Chen, C.-H.; Carta, V.; Moënne-Loccoz, P.; Smith, J. M. ACS Cent. Sci. 2020, 6, 1572–1577. Valdez-Moreira, J. A.; Thorarinsdottir, A. E.; DeGayner, J. A.; Lutz, S. A.; Chen, C.-H.; Losovyj, Y.; Pink, M.; Harris, T. D.; Smith, J. M. J. Am. Chem. Soc. 2019, 141, 17092–17097. Martinez, J. L.; Lutz, S. A.; Yang, H.; Xie, J.; Telser, J.; Hoffman, B. M.; Carta, V.; Pink, M.; Losovyj, Y.; Smith, J. M. Science 2020, 370, 356. Hickey, A. K.; Greer, S. M.; Valdez-Moreira, J. A.; Lutz, S. A.; Pink, M.; DeGayner, J. A.; Harris, T. D.; Hill, S.; Telser, J.; Smith, J. M. J. Am. Chem. Soc. 2019, 141, 11970–11975. Lutz, S. A.; Hickey, A. K.; Gao, Y.; Chen, C.-H.; Smith, J. M. J. Am. Chem. Soc. 2020, 142, 15527–15535. Forshaw, A. P.; Bontchev, R. P.; Smith, J. M. Inorg. Chem. 2007, 46, 3792–3794. Scepaniak, J. J.; Margarit, C. G.; Bontchev, R. P.; Smith, J. M. Acta Crystallogr. C 2013, 69, 968–971. Forshaw, A. P.; Smith, J. M.; Ozarowski, A.; Krzystek, J.; Smirnov, D.; Zvyagin, S. A.; Harris, T. D.; Karunadasa, H. I.; Zadrozny, J. M.; Schnegg, A.; Holldack, K.; Jackson, T. A.; Alamiri, A.; Barnes, D. M.; Telser, J. Inorg. Chem. 2013, 52, 144–159. Colmer, H. E.; Margarit, C. G.; Smith, J. M.; Jackson, T. A.; Telser, J. Eur. J. Inorg. Chem. 2016, 2016, 2413–2423. Ding, M.; Cutsail Iii, G. E.; Aravena, D.; Amoza, M.; Rouzières, M.; Dechambenoit, P.; Losovyj, Y.; Pink, M.; Ruiz, E.; Clérac, R.; Smith, J. M. Chem. Sci. 2016, 7, 6132–6140. Cowley, R. E.; Bontchev, R. P.; Duesler, E. N.; Smith, J. M. Inorg. Chem. 2006, 45, 9771–9779. Cowley, R. E.; Bontchev, R. P.; Sorrell, J.; Sarracino, O.; Feng, Y.; Wang, H.; Smith, J. M. J. Am. Chem. Soc. 2007, 129, 2424–2425. Xu, S.; Boschen, J. S.; Biswas, A.; Kobayashi, T.; Pruski, M.; Windus, T. L.; Sadow, A. D. Dalton Trans. 2015, 44, 15897–15904. Xu, S.; Manna, K.; Ellern, A.; Sadow, A. D. Organometallics 2014, 33, 6840–6860. Xu, S.; Everett, W. C.; Ellern, A.; Windus, T. L.; Sadow, A. D. Dalton Trans. 2014, 43, 14368–14376. Spicer, M. D.; Reglinski, J. Eur. J. Inorg. Chem. 2009, 2009, 1553–1574. Neshat, A.; Afrasi, M.; Gilanchi, S.; Gholinejad, M. ChemistrySelect 2019, 4, 9268–9273. Tran, B. L.; Carrano, C. J. Inorg. Chem. 2007, 46, 5429–5438. Hill, M.; Kehr, G.; Erker, G.; Kataeva, O.; Fröhlich, R. Chem. Commun. 2004. https://doi.org/10.1039/B400228H, 1020-1021. Hill, M.; Erker, G.; Kehr, G.; Fröhlich, R.; Kataeva, O. J. Am. Chem. Soc. 2004, 126, 11046–11057. Sun, Y.; Spence, R. E. V. H.; Piers, W. E.; Parvez, M.; Yap, G. P. A. J. Am. Chem. Soc. 1997, 119, 5132–5143. Sun, Y.; Piers, W. E.; Yap, G. P. A. Organometallics 1997, 16, 2509–2513. Shapiro, P. J. Eur. J. Inorg. Chem. 2001, 2001, 321–326. Larkin, S. A.; Golden, J. T.; Shapiro, P. J.; Yap, G. P. A.; Foo, D. M. J.; Rheingold, A. L. Organometallics 1996, 15, 2393–2398. Stelck, D. S.; Shapiro, P. J.; Basickes, N.; Rheingold, A. L. Organometallics 1997, 16, 4546–4550. Shapiro, P. J.; Jiang, F.; Jin, X.; Twamley, B.; Patton, J. T.; Rheingold, A. L. Eur. J. Inorg. Chem. 2004, 2004, 3370–3378. Reetz, M. T.; Willuhn, M.; Psiorz, C.; Goddard, R. Chem. Commun. 1999. https://doi.org/10.1039/A902543J, 1105-1106. Lancaster, S. J.; Bochmann, M. Organometallics 2001, 20, 2093–2101. Aldridge, S.; Bresner, C. Coord. Chem. Rev. 2003, 244, 71–92. Manna, K.; Ellern, A.; Sadow, A. D. Chem. Commun. 2010, 46, 339–341. Manna, K.; Xu, S.; Sadow, A. D. Angew. Chem. Int. Ed. 2011, 50, 1865–1868. Manna, K.; Everett, W. C.; Schoendorff, G.; Ellern, A.; Windus, T. L.; Sadow, A. D. J. Am. Chem. Soc. 2013, 135, 7235–7250. Manna, K.; Eedugurala, N.; Sadow, A. D. J. Am. Chem. Soc. 2015, 137, 425–435. Manna, K.; Kruse, M. L.; Sadow, A. D. ACS Catalysis 2011, 1, 1637–1642. Kanbur, U.; Ellern, A.; Sadow, A. D. Organometallics 2018, 37, 4409–4414. Kanbur, U.; Sadow, A. D. Chem. A Eur. J. 2020, 26, 5479–5493. Boardman, B. M.; Bazan, G. C. Acc. Chem. Res. 2009, 42, 1597–1606. Lee, B. Y.; Bazan, G. C.; Vela, J.; Komon, Z. J. A.; Bu, X. J. Am. Chem. Soc. 2001, 123, 5352–5353. Carone, C. L. P.; Bisatto, R.; Galland, G. B.; Rojas, R.; Bazan, G. J. Polym. Sci. A Polym. Chem. 2008, 46, 54–59. Coffin, R. C.; Schneider, Y.; Kramer, E. J.; Bazan, G. C. J. Am. Chem. Soc. 2010, 132, 13869–13878. Chen, Y.; Boardman, B. M.; Wu, G.; Bazan, G. C. J. Organomet. Chem. 2007, 692, 4745–4749. Lee, B. Y.; Bu, X.; Bazan, G. C. Organometallics 2001, 20, 5425–5431. Kim, Y. H.; Kim, T. H.; Lee, B. Y.; Woodmansee, D.; Bu, X.; Bazan, G. C. Organometallics 2002, 21, 3082–3084. Chen, Y.; Wu, G.; Bazan, G. C. Angew. Chem. Int. Ed. 2005, 44, 1108–1112. Bonnet, M. C.; Dahan, F.; Ecke, A.; Keim, W.; Schulz, R. P.; Tkatchenko, I. J. Chem. Soc. Chem. Commun. 1994. https://doi.org/10.1039/C39940000615, 615-616. Komon, Z. J. A.; Bu, X.; Bazan, G. C. J. Am. Chem. Soc. 2000, 122, 1830–1831. Komon, Z. J. A.; Bu, X.; Bazan, G. C. J. Am. Chem. Soc. 2000, 122, 12379–12380. Contrella, N. D.; Jordan, R. F. Organometallics 2014, 33, 7199–7208. Wilders, A. M.; Contrella, N. D.; Sampson, J. R.; Zheng, M.; Jordan, R. F. Organometallics 2017, 36, 4990–5002. Shim, C. B.; Kim, Y. H.; Lee, B. Y.; Dong, Y.; Yun, H. Organometallics 2003, 22, 4272–4280. Shim, C. B.; Kim, Y. H.; Lee, B. Y.; Shin, D. M.; Chung, Y. K. J. Organomet. Chem. 2003, 675, 72–76. Kwon, H. Y.; Lee, S. Y.; Lee, B. Y.; Shin, D. M.; Chung, Y. K. Dalton Trans. 2004. https://doi.org/10.1039/B317033K, 921-928. Thomas, C. M.; Peters, J. C. Inorg. Chem. 2004, 43, 8–10. Beach, M. T.; Walker, J. M.; Larocque, T. G.; Deagle, J. L.; Wang, R.; Spivak, G. J. J. Organomet. Chem. 2008, 693, 2921–2928. Granville, S. L.; Welch, G. C.; Stephan, D. W. Inorg. Chem. 2012, 51, 4711–4721. Jaska, C. A.; Dorn, H.; Lough, A. J.; Manners, I. Chem. A Eur. J. 2003, 9, 271–281.

420 241. 242. 243. 244. 245. 246. 247. 248. 249. 250. 251. 252. 253. 254. 255. 256. 257. 258. 259. 260. 261. 262. 263. 264. 265. 266. 267. 268. 269. 270. 271. 272. 273. 274. 275. 276. 277. 278. 279. 280. 281. 282. 283. 284. 285. 286. 287. 288. 289. 290. 291. 292. 293. 294. 295. 296. 297. 298. 299. 300. 301. 302. 303. 304.

Ligands Featuring Covalently Tethered Moderate to Weakly Coordinating Anions Bontemps, S.; Bouhadir, G.; Miqueu, K.; Bourissou, D. J. Am. Chem. Soc. 2006, 128, 12056–12057. Fischbach, A.; Bazinet, P. R.; Waterman, R.; Tilley, T. D. Organometallics 2008, 27, 1135–1139. Konishi, S.; Iwai, T.; Sawamura, M. Organometallics 2018, 37, 1876–1883. Gott, A. L.; Piers, W. E.; Dutton, J. L.; McDonald, R.; Parvez, M. Organometallics 2011, 30, 4236–4249. Kim, Y.; Jordan, R. F. Organometallics 2011, 30, 4250–4256. Qiao, S.; Hoic, D. A.; Fu, G. C. J. Am. Chem. Soc. 1996, 118, 6329–6330. Hoic, D. A.; Davis, W. M.; Fu, G. C. J. Am. Chem. Soc. 1996, 118, 8176–8177. Holschumacher, D.; Bannenberg, T.; Hrib, C. G.; Jones, P. G.; Tamm, M. Angew. Chem. Int. Ed. 2008, 47, 7428–7432. Kronig, S.; Theuergarten, E.; Daniliuc, C. G.; Jones, P. G.; Tamm, M. Angew. Chem. Int. Ed. 2012, 51, 3240–3244. Wang, Y.; Xie, Y.; Abraham, M. Y.; Wei, P.; Schaefer, H. F.; Schleyer, P. V. R.; Robinson, G. H. J. Am. Chem. Soc. 2010, 132, 14370–14372. Kolychev, E. L.; Kronig, S.; Brandhorst, K.; Freytag, M.; Jones, P. G.; Tamm, M. J. Am. Chem. Soc. 2013, 135, 12448–12459. Koneczny, M.; Phong Ho, L.; Nasr, A.; Freytag, M.; Jones, P. G.; Tamm, M. Adv. Synth. Catal. 2020, 362, 3857–3863. Winkler, A.; Brandhorst, K.; Freytag, M.; Jones, P. G.; Tamm, M. Organometallics 2016, 35, 1160–1169. Frosch, J.; Freytag, M.; Jones, P. G.; Tamm, M. J. Organomet. Chem. 2020, 918, 121311. Igarashi, A.; Kolychev, E. L.; Tamm, M.; Nomura, K. Organometallics 2016, 35, 1778–1784. Nagai, G.; Mitsudome, T.; Tsutsumi, K.; Sueki, S.; Ina, T.; Tamm, M.; Nomura, K. J. Jpn. Pet. Inst. 2017, 60, 256–262. Nomura, K.; Nagai, G.; Izawa, I.; Mitsudome, T.; Tamm, M.; Yamazoe, S. ACS Omega 2019, 4, 18833–18845. Nomura, K.; Nagai, G.; Nasr, A.; Tsutsumi, K.; Kawamoto, Y.; Koide, K.; Tamm, M. Organometallics 2019, 38, 3233–3244. Phillips, N.; Tirfoin, R.; Aldridge, S. Dalton Trans. 2014, 43, 15279–15282. Phillips, N.; Dodson, T.; Tirfoin, R.; Bates, J. I.; Aldridge, S. Chem. A Eur. J. 2014, 20, 16721–16731. Jana, A.; Azhakar, R.; Tavcar, G.; Roesky, H. W.; Objartel, I.; Stalke, D. Eur. J. Inorg. Chem. 2011, 2011, 3686–3689. Pranckevicius, C.; Stephan, D. W. Chem. A Eur. J. 2014, 20, 6597–6602. Wacker, A.; Yan, C. G.; Kaltenpoth, G.; Ginsberg, A.; Arif, A. M.; Ernst, R. D.; Pritzkow, H.; Siebert, W. J. Organomet. Chem. 2002, 641, 195–202. Nasr, A.; Winkler, A.; Tamm, M. Coord. Chem. Rev. 2016, 316, 68–124. Weiss, A.; Pritzkow, H.; Siebert, W. Eur. J. Inorg. Chem. 2002, 2002, 1607–1614. Pellei, M.; Vallesi, R.; Bagnarelli, L.; Dias, H. V. R.; Santini, C. Molecules 2020, 25. Eleazer, B. J.; Peryshkov, D. V. Comments Inorg. Chem. 2018, 38, 79–109. Kirchmann, M.; Wesemann, L. Dalton Trans. 2008. https://doi.org/10.1039/B715305H, 444-446. Hertler, W. R.; Raasch, M. S. J. Am. Chem. Soc. 1964, 86, 3661–3668. Kirchmann, M.; Wesemann, L. Dalton Trans. 2008. https://doi.org/10.1039/B718569C, 2144-2148. Zhang, Y.; Sun, Y.; Lin, F.; Liu, J.; Duttwyler, S. Angew. Chem. Int. Ed. 2016, 55, 15609–15614. Zhang, Y.; Wang, T.; Wang, L.; Sun, Y.; Lin, F.; Liu, J.; Duttwyler, S. Chem. A Eur. J. 2018, 24, 15812–15817. Jasper, S. A.; Jones, R. B.; Mattern, J.; Huffman, J. C.; Todd, L. J. Inorg. Chem. 1994, 33, 5620–5624. Jasper, S. A.; Mattern, J.; Huffman, J. C.; Todd, L. J. Polyhedron 2007, 26, 3793–3798. Bregadze, V. I.; Sivaev, I. B.; Gabel, D.; WöHrle, D. J. Porphyrins Phthalocyanines 2001, 05, 767–781. Semioshkin, A.; Tsaryova, O.; Zhidkova, O.; Bregadze, V.; Wöhrle, D. J. Porphyrins Phthalocyanines 2006, 10, 1293–1300. Semioshkin, A. A.; Sivaev, I. B.; Bregadze, V. I. Dalton Trans. 2008. https://doi.org/10.1039/B715363E, 977-992. El-Zaria, M. E.; Ban, H. S.; Nakamura, H. Chem. A Eur. J. 2010, 16, 1543–1552. Nedunchezhian, K.; Aswath, N.; Thiruppathy, M.; Thirugnanamurthy, S. J. Clin. Diagn. Res. 2016, 10, ZE01–ZE04. Birsöz, B.; Nar, I.; Gül, A. J. Organomet. Chem. 2014, 755, 64–71. Voloshin, Y. Z.; Varzatskii, O. A.; Zhizhin, K. Y.; Kuznetsov, N. T.; Bubnov, Y. N. Russ. Chem. Bull. 2006, 55, 22–25. Peymann, T.; Knobler, C. B.; Hawthorne, M. F. Inorg. Chem. 2000, 39, 1163–1170. Himmelspach, A.; Finze, M.; Raub, S. Angew. Chem. Int. Ed. 2011, 50, 2628–2631. Schäfer, M.; Krummenacher, I.; Braunschweig, H.; Finze, M. Z. Anorg. Allg. Chem. 2015, 641, 660–668. Hailmann, M.; Wolf, N.; Renner, R.; Hupp, B.; Steffen, A.; Finze, M. Chem. A Eur. J. 2017, 23, 11684–11693. Hailmann, M.; Wolf, N.; Renner, R.; Schäfer, T. C.; Hupp, B.; Steffen, A.; Finze, M. Angew. Chem. Int. Ed. 2016, 55, 10507–10511. Röhrscheid, F.; Holm, R. H. J. Organomet. Chem. 1965, 4, 335–338. El-Hellani, A.; Kefalidis, C. E.; Tham, F. S.; Maron, L.; Lavallo, V. Organometallics 2013, 32, 6887–6890. Molinos, E.; Kociok-Köhn, G.; Weller, A. S. Chem. Commun. 2005. https://doi.org/10.1039/B504630K, 3609-3611. Rifat, A.; Patmore, N. J.; Mahon, M. F.; Weller, A. S. Organometallics 2002, 21, 2856–2865. Rifat, A.; Kociok-Köhn, G.; Steed, J. W.; Weller, A. S. Organometallics 2003, 23, 428–432. Weller, A. S.; Mahon, M. F.; Steed, J. W. J. Organomet. Chem. 2000, 614-615, 113–119. Molinos, E.; Brayshaw, S. K.; Kociok-Kohn, G.; Weller, A. S. Dalton Trans. 2007. https://doi.org/10.1039/B711468K, 4829-4844. Molinos, E.; Brayshaw, S. K.; Kociok-Köhn, G.; Weller, A. S. Organometallics 2007, 26, 2370–2382. Mhinzi, G. S.; Litster, S. A.; Redhouse, A. D.; Spencer, J. L. J. Chem. Soc. Dalton Trans. 1991. https://doi.org/10.1039/DT9910002769, 2769-2776. Lavallo, V.; Wright, J. H., II; Tham, F. S.; Quinlivan, S. Angew. Chem. Int. Ed. 2013, 52, 3172–3176. Estrada, J.; Woen, D. H.; Tham, F. S.; Miyake, G. M.; Lavallo, V. Inorg. Chem. 2015, 54, 5142–5144. Chan, A. L.; Estrada, J.; Kefalidis, C. E.; Lavallo, V. Organometallics 2016, 35, 3257–3260. Kleinsasser, J. F.; Reinhart, E. D.; Estrada, J.; Jordan, R. F.; Lavallo, V. Organometallics 2018, 37, 4773–4783. Jones, C. G.; Asay, M.; Kim, L. J.; Kleinsasser, J. F.; Saha, A.; Fulton, T. J.; Berkley, K. R.; Cascio, D.; Malyutin, A. G.; Conley, M. P.; Stoltz, B. M.; Lavallo, V.; Rodríguez, J. A.; Nelson, H. M. ACS Cent. Sci. 2019, 5, 1507–1513. El-Hellani, A.; Lavallo, V. Angew. Chem. Int. Ed. 2014, 53, 4489–4493. Asay, M. J.; Fisher, S. P.; Lee, S. E.; Tham, F. S.; Borchardt, D.; Lavallo, V. Chem. Commun. 2015, 51, 5359–5362. Fisher, S. P.; El-Hellani, A.; Tham, F. S.; Lavallo, V. Dalton Trans. 2016, 45, 9762–9765. Fisher, S. P.; McArthur, S. G.; Tej, V.; Lee, S. E.; Chan, A. L.; Banda, I.; Gregory, A.; Berkley, K.; Tsay, C.; Rheingold, A. L.; Guisado-Barrios, G.; Lavallo, V. J. Am. Chem. Soc. 2020, 142, 251–256.

1.13

Redox-Active Ligands in Organometallic Chemistry

Errikos Kounalis and Daniël LJ Broere, Organic Chemistry and Catalysis, Debye Institute for Nanomaterials Science Faculty of Science, Utrecht University, Universiteitsweg 99, Utrecht, The Netherlands © 2022 Elsevier Ltd. All rights reserved.

1.13.1 Introduction 1.13.2 Redox-active ligands as electron reservoirs 1.13.2.1 Early transition metals 1.13.2.2 Late transition metals 1.13.2.3 Other metals 1.13.3 Modification of the Lewis acid-base properties through ligand-centered redox changes 1.13.4 Redox-active ligand-to-substrate single-electron transfer 1.13.5 Cooperative ligand-centered reactivity 1.13.6 Conclusion Acknowledgment References

1.13.1

421 423 423 425 430 431 432 435 439 439 439

Introduction

The assignment of formal oxidation states of metals is a core concept that has its clearest applicability for the balancing of redox reactions. In addition, the assignment of formal metal oxidation states is often used to predict or rationalize observed spectroscopic features, magnetic behavior and chemical reactivity, but can encounter problems.1 A much more pronounced limitation of the oxidation state formalism is its inability to account for ligand-centered radicals.2 This was first brought to light in a series of reports from the 1960s involving complexes of 1,2-dithiolene ligands.3 Although it is currently well understood that these ligands can bind metals in a di-anionic ene-1,2-dithiolate, a mono-anionic radical and a neutral 1,2-dithioketone form (Fig. 1),4 these ligands sparked numerous discussions. This perhaps is most evident in a seminal review on these spectroscopic curiosities by Jørgensen,5 who therein categorized sulfur-containing ligands as “innocent” or “suspect” based on whether the ligand allowed for the metal’s formal oxidation state to be defined or not, respectively. Nearly 50 years after Jørgensen’s seminal report, “suspect” ligands have become better known as “noninnocent” or “redoxactive.” Although these terms are often used interchangeably, the former implies an uncertainty in the assignment of oxidation states.6 In contrast, the term “redox-active” is appropriate when observed redox-changes are (partially) ligand-centered, and the metal’s d-electronic configuration can be experimentally determined or derived as a physical or spectroscopic oxidation state.7 This terminology is no longer limited to the 1,2-dithiolene ligands, but has been applied to a large variety of ligand classes, including diarylamines,8 diimines,9 iminopyridines,10 catecholates,11 o-phenylene diamines,12 o-aminophenols,13 verdazyls,14 formazanates,15 bipyridines,16 aminotroponiminates,17 arylazopyridines18 and several others.19 Moreover, the field has evolved far beyond the determination of oxidation states as these ligands have been shown to expand upon a metal’s common reactivity. As the focus of this chapter is on the latter, the term “redox-active” will be used throughout, despite that in some cases there might be ambiguity about oxidation states. In typical metal complexes bearing “redox-inactive” ligands, the energy required for ligand-centered redox events is substantially higher than that for redox events (oxidation state changes) that are localized on the metal center. Hence, changes in electronic structure associated with chemical transformations, using in this example an oxidative addition, typically occur at the metal and not at the coordinated ligands (Fig. 2, top). In metal complexes bearing redox-active ligands (RALs), the energy required for ligand-centered redox changes can be lower than – or similar to – what is required for metal-centered reduction or oxidation. Such metal complexes commonly feature a p-symmetric, ligand-centered molecular orbital (MO) that is higher or similar in energy than the MOs that are predominantly localized on the metal atom. Consequently, in an oxidative addition reaction either solely ligand-centered redox processes can occur, with the metal center remaining in the same oxidation state (Fig. 2, middle), or both the ligand and metal change oxidation state in a synergistic fashion (Fig. 2, bottom).20 Likewise, a low-energy, unfilled p-symmetric ligand-centered MO can accept one or two of the electrons in a reductive elimination. This interesting feature of RALs enables them to “expand” upon a metal’s common palette of oxidation states by providing the necessary electrons, avoiding unfavorable oxidation states or enabling two-electron transformations on metals that are more prone toward one-electron transformations. Transition metals have found widespread application in catalysis due to their ability to adopt multiple oxidation states. The electronic configuration and preference for one- or two-electron redox changes of each metal has resulted in the application of certain groups of metals for specific chemical transformations. For example, the ability of Pd and Pt to shuttle between oxidation states by two-electron increments during metal-mediated bond making/breaking steps has enabled their extensive application in cross-coupling reactions.21 By tuning the steric and donating properties of classical ancillary ligands, the activity, selectivity or stability of transition metal catalysts can be greatly enhanced.22 However, such ligand modifications typically do not provide access to new energetically accessible ligand-centered energy levels, which enable new chemical transformation beyond the intrinsic reactivity of a metal. In addition to the search for new chemical transformations, there is a growing impetus to replace existing

Comprehensive Organometallic Chemistry IV

https://doi.org/10.1016/B978-0-12-820206-7.00028-7

421

Redox-Active Ligands in Organometallic Chemistry

422

Fig. 1 The three limiting structures that are possible for a neutral metal complex bearing two 1,2-dithiolene ligands.

Fig. 2 Various possibilities for oxidation state changes in an oxidative addition reaction. Top: A “classical” redox-inactive ligand where the metal is oxidized by two electrons. Middle: Redox-active ligand functioning as a two-electron reservoir, thereby keeping the metal in the same oxidation state. Bottom: synergistic redox changes on both the metal and redox-active ligand.

catalysts based on scarce metals with analogs based on more abundant metals. Redox-active ligands have been shown to be able to contribute to both these goals through enabling reactivity that is not typically associated with certain metals or by facilitating two-electron redox changes on metals that are prone to undergo one-electron changes.23 Nature employs redox-active moieties in various metalloenzyme active sites,24 wherein the multi-electron reactions are broken down into smaller steps, enabling reactions to occur near thermodynamic potential (see Section 1.13.5).25 Inspired by natural RALs in metalloenzyme active sites, chemists have discovered that RALs can be utilized to expand upon a metal’s “normal” reactivity, thereby offering a new tool to tackle current challenges in chemistry.26 In this chapter we will describe the four main strategies how RALs can enable new chemical reactivity (Fig. 3). The examples described in the following sections are by no means comprehensive, (A)

(B)

(C)

(D)

Fig. 3 The four main strategies how redox-active ligands can expand upon a metal’s reactivity.

Redox-Active Ligands in Organometallic Chemistry

423

but are selected to demonstrate the diversity in which RALs can be used. The main strategy used involves RALs to influence chemical reactivity is by (A) acting as an electron reservoir to enable two-electron transformations. Ligand-centered redox changes can also be used to (B) modify the Lewis acidity of the bound metal and bound substrates. RALs can also be utilized to (C) transfer a single electron toward a redox-active substrate, thereby generating a reactive ligand-centered radical.27 The final strategy that will be discussed involves (D) the utilization of reactive ligand-centered radicals that are actively involved in chemical transformations through the reversible making/breaking of covalent bonds. We will focus on the strategies where these ligands act as a supporting ligand. Monoatomic ligands such as oxos and nitridos are not covered. Similarly, single-electron transfer from metals to form carbene and nitrene centered radicals28 will be only briefly discussed in the context of strategy (C).

1.13.2

Redox-active ligands as electron reservoirs

1.13.2.1

Early transition metals

Oxidative addition and reductive elimination are key reactions in transition metal catalysis and are most commonly associated with late transition metals. These reactions are classically associated with a decrease or increase of the dn electronic count by two electrons, respectively. Hence, a metal complex with a d0 configuration, such as Zr(IV), should not be able to perform an oxidative addition reaction. Similarly, a reductive elimination from such a complex would result in an unfavorable Zr(II) oxidation state. By creating more accessible ligand-centered energy levels for redox changes, RALs can function as electron reservoirs that can provide or accept the electrons associated with a chemical transformation. Consequently, RALs can enable classical elementary steps that involve two-electron redox changes in high-valent early transition metal complexes, thereby expanding upon the classical reactivity paradigm. An early example of this concept was reported by Heyduk and co-workers, who demonstrated that redox-active aminophenol-based ligands can enable the oxidative addition of Cl2 to a Zr(IV) complex featuring two redox-active o-amidophenolate ligands (Scheme 1, top).12a This formal oxidative addition reaction is facilitated by one-electron oxidation of each o-amidophenolate ligand, which are converted to the iminosemiquinonato state. The one-electron oxidation of the ligands was confirmed by UV-Vis and EPR spectroscopy and through analysis of the bond metrics29 in the solid-state structure obtained by single-crystal X-ray diffraction. Although the reverse CldCl reductive elimination proved to be inaccessible, an in-situ prepared Zr(IV) diphenyl analog (Scheme 1, bottom) was found to reductively eliminate biphenyl concomitant with reduction of the two iminosemiquinonato ligands to their amidophenolate state, thereby evading the unfavorable Zr(II) state.30 Similar RAL-mediated oxidative addition reactivity on Zr(IV) was reported by the same group using o-phenylenediamido ligands.31

Scheme 1 The oxidative addition (top) and reductive elimination (bottom) from Zr(IV) complexes enabled by redox-active ligands functioning as an electron reservoir.

The more oxidized states of most RALs are typically neutral or mono-anionic. This can often result in weaker binding properties, especially when dealing with high oxidation state early transition metals. Hence, their incorporation into scaffolds of higher denticity can prevent detrimental ligand dissociation in catalytic reactions. Along these lines, the Heyduk group synthesized a Zr(IV) complex – featuring a tetradentate RAL previously developed by the Wieghardt group32 – capable of catalytic diphenylhydrazine disproportionation.33 The proposed mechanism of this catalytic reaction is depicted in Scheme 2 and starts with dissociation of a THF ligand (A), enabling binding of diphenylhydrazine (B). Subsequent 1,2 N-H elimination of aniline leads to the formation of the corresponding imido ligand (C), and the required electrons for this step are provided by the bis(amidophenolate) scaffold, which undergoes a two-electron oxidation. Subsequently, the Zr imido abstracts two H-atoms from a diphenylhydrazine molecule to give azobenzene, and this is concomitant with two electron reduction of the RAL (D), which regains its tetra-anionic form. After aniline dissociation a vacant site is regenerated completing the cycle (E).

424

Redox-Active Ligands in Organometallic Chemistry

Scheme 2 Catalytic diphenylhydrazine disproportionation by a Zr(IV) complex bearing a tetradentate redox-active ligand.

Besides the tetradentate bis(amidophenolate) ligand, the Heyduk group also explored a family of tridentate “pincer-type” NNN and ONO ligands for catalytic nitrene transfer reactions.34 A Ta(V) complex bearing a tri-anionic redox-active ONO ligand and two tert-butoxide ligands was found to catalyze the conversion of phenylazide to the corresponding azobenzene and dinitrogen (Scheme 3), albeit with only 5 turnovers after 7 days at 55 oC.33 Unlike the analog bearing two chloride ligands instead of alkoxides, the catalyst did not decompose in the excess of phenylazide. The proposed mechanism starts with a reaction of (ONOred)Ta(Ot-Bu)2 with PhN3 to form an imido complex concomitant with N2 formation (A). The two electrons required for

Scheme 3 Catalytic azobenzene synthesis from aryl azides through a four-electron reductive elimination by a Ta(V) complex bearing a redox-active ONO ligand.

Redox-Active Ligands in Organometallic Chemistry

425

this transformation are provided by two-electron oxidation of the tridentate RAL. A subsequent dimerization (B) and overall four-electron reductive elimination (C) gives azobenzene and completes the cycle. Another interesting example of how two RALs can enable multi-electron processes was reported by Abu-Omar and co-workers in 2007.35 They demonstrated that a Zr(IV) complex bearing two RALs, which are best described as diamides that also donate electron density from the olefinic p-bond (Scheme 4), reacts with anhydrous O2 to give an unprecedented bis-peroxo-Zr(IV) complex. The solid-state structure showed that the peroxides bind in Z2 fashion to the Zr(IV) center and that both bidentate amine ligands underwent a two-electron oxidation, returning to their diimine form. This example demonstrates that RALs can enable an overall four-electron reduction of two O2 molecules on Zr(IV) without changing metal oxidation state.

Scheme 4 Multielectron activation of O2 to give a bisperoxo Zr complex.

Related redox-active glyoxal diamine ligands have also been employed by the Mashima group in Ta36 and Nb37 complexes for the reductive cleavage of CdCl bonds in haloalkanes and catalytic radical addition reactions. Using a diimino-trichloro-niobium compound, they demonstrated the catalytic atom transfer radical addition (ATRA) of CdCl bonds on terminal and cyclic olefins (Scheme 5), with the diimine ligand playing a crucial role in storing and releasing the electron necessary for the transformation. The proposed mechanism starts from the off-cycle NbCl3-styrene complex, which undergoes ligand-centered reduction upon styrene dissociation (A). Subsequently, the RAL is reoxidized to the iminoamidate state upon precoordination of the C-bound Cl to the Nb center (B), before it is abstracted (C). The formed organic radical fragment then adds to the double bond of styrene (D). Subsequently, the resulting organic radical fragment re-abstracts the Cl atom bound to Nb, to reform the active species upon release of the organic product (E). Although there is no clear advantage for using the reported system in favor of well-established ATRA catalysts,38 it is an interesting distinct example on how RALs can expand upon a metal’s common reactivity.

Scheme 5 Catalytic radical addition of CdCl bonds to alkenes mediated by reversible ligand-centered redox changes.

1.13.2.2

Late transition metals

One of the key challenges in modern chemistry is the replacement of catalysts based on scarce noble metals (such as Pd, Pt, Rh) with alternatives based on earth-abundant metals such as Fe, Co and Ni.39,40 A major obstacle in this pursuit is the difference in electronic structure between these metals. The more abundant first-row transition metals commonly undergo more facile one-electron redox changes. This can be problematic for bond making and breaking events that rely on two-electron redox changes with noble metals.

426

Redox-Active Ligands in Organometallic Chemistry

RALs are capable of conferring “nobility” on first-row transition metals by providing additional energetically accessible levels for redox changes.23 This can enable a metal to maintain its electronic configuration as all oxidation state changes occur on the ligand. Alternatively, the RAL and metal can work in synergy by each undergoing a single-electron oxidation or reduction. Pyridinediimine (PDI) ligands were initially explored by the groups of Gibson41,42 and Brookhart43 in the late 90’s, for their use in iron- and cobalt-mediated olefin oligo- and polymerization. The redox-active behavior of this ligand was established relatively early by the Budzelaar and Wieghardt groups. They demonstrated that for bis-ligated, octahedral, divalent base metal complexes, these ligands can either be bound in a neutral, radical mono-anionic or diradical di-anionic form (Scheme 6).44

R R

R

N N

N [PDI]

0

e-e

R

R

-

R

R

N N

N [PDI]

e-e

R

R

-

R

R

N N

-

N [PDI]

R

2-

Scheme 6 The various oxidation states of the pyridine diimine (PDI) ligand.

A few years later, the Chirik group used the iPrPDI ligand to synthesize the first example of an iron(0) bis(dinitrogen) complex (Scheme 7).45 In solution, a N2 molecule is liberated, forming the tetracoordinate complex. These and related complexes have been investigated using Mössbauer spectroscopy, (TD)-DFT, XRD, XAS and XES in great detail46 and it is believed that the mono(dinitrogen) complex is best described as an intermediate spin Fe(II) complex, weakly antiferromagnetically coupled to the iPr PDI ligand in its diradical, di-anionic triplet state (Scheme 6, left). The bis(dinitrogen) complex, however, is better described as a highly covalent compound, existing as a resonance structure between the extremes of an Fe(0) center bound to a neutral PDI ligand and a Fe(II) ion bound to the PDI ligand in its diradical, di-anionic form (Scheme 6, right).

Scheme 7 The proposed electronic structure of [iPrPDI]Fe(N2) and the two resonance forms of [iPrPDI]Fe(N2)2.

The bis(dinitrogen) complex compound has proven to be an excellent pre-catalyst for a variety of transformations,47 including selective non-Markovnikov hydrosilylations of olefins and alkynes, hydrogenations of alkynes and olefins under mild conditions and [2 +2] cycloaddition of a-o dienes yielding bicyclic systems containing a cyclobutane ring.48 The authors initially proposed that throughout the catalytic cycle for the latter reaction the iron center remains in a ferrous state, and involves shuttling of the PDI ligand between the diradical di-anionic and neutral states. However, high level computational studies, including multi-reference methods, by Hu and Chen suggested a substrate-dependent two-state reactivity, and that the RAL renders a thermodynamically more favorable Fe(I)/Fe(III) cycle in CdC coupling.49 Extensive experimental and computational follow-up studies by Chirik and coworkers also supported a substrate dependent mechanism. For the intramolecular [2+ 2] cycloaddition of a-o dienes (Scheme 8) the metallacyclic intermediate was found to be the resting state, and the catalytic cycle was proposed to remain on the S ¼ 1 spin surface involving a Fe(I)/Fe(III) cycle. In contrast, for the intermolecular [2 +2] cycloaddition of a-olefins an S ¼ 0 iron complex

Scheme 8 The proposed mechanism for the catalytic intramolecular [2+2] cycloaddition of a-o dienes wherein [iPrPDI]Fe(N2)2 serves as a pre-catalyst.

Redox-Active Ligands in Organometallic Chemistry

427

bearing a N2 and olefin ligand was found to be the resting state. The proposed mechanism for this reaction (Scheme 9) involves a spin transition upon reductive elimination from the S ¼ 1 metallacycle to give an S ¼ 0 dinitrogen complex featuring a different PDI oxidation state. Although the authors state that it is ambiguous whether the observed resting state is on- or off cycle, they only observed a negative order in [N2] for the intermolecular reaction. This change in mechanism is proposed to be due to the chelate effect, which makes a diene bind more strongly than N2.50

Scheme 9 The proposed mechanism for the catalytic intermolecular [2 +2] cycloaddition of a-olefins in which [iPrPDI]Fe(N2)2 could be an on-cycle intermediate.

Expanding on the PDI system, the Uyeda group developed the naphthyridine diimine ligand (NDI) that can host two metal centers.51 This RAL can bind two metals in close proximity, which enables metal-metal cooperativity, opening up more avenues in its reactivity. Combining these two types of reactivity for Ni complexes, the group has reported dinuclear oxidative addition,52 carbene transfers,53 carbonylations,54 and azoarene synthesis.55 More recently, they reported that this dinuclear complex bearing a RAL can catalyze the reductive [4 +1] cycloadditions of vinylidenes and dienes.56 The latter is proposed to proceed through the catalytic cycle depicted in Scheme 10. In the first step the 1,1-dichloroalkene undergoes oxidative addition to one of the Ni centers, and this step is facilitated by one-electron oxidation of the NDI ligand (A). Subsequent reduction by Zn results in the formation of a bridging vinylidene species through oxidative addition of the remaining CdCl bond (B). After binding a 1,3-diene (C) a migratory

Scheme 10 Catalytic reductive [4 +1] cycloaddition of dienes with vinylidenes mediated by a dinickel complex bearing a redox-active NDI ligand.

428

Redox-Active Ligands in Organometallic Chemistry

insertion forms a nickelacycle (D). In the final step the [4 + 1] cycloadduct is released by a reductive elimination that also results in reduction of the redox-active NDI ligand. Although the NDI ligand has been demonstrated to be redox-active and is proposed to actively take part in redox changes throughout the proposed catalytic cycle, its role has not been investigated to the same extent as for the mononuclear Fe PDI systems described above. Although the redox-active nature of formazans has been exploited in colorimetric assays for assessing cell metabolic activity since the 1980s,57 their potential as RALs has not been recognized until recently. In 2014, these nitrogen rich analogs of b-diketiminates were shown by Otten and co-workers to serve as an electron reservoir in tetrahedral Zn complexes bearing two mono-anionic formazanate ligands.58 Treatment of this homoleptic Zn(II) complex with one or two equivalents of sodium amalgam provides access to complexes bearing one or two ligand radicals, respectively (Scheme 11).

Scheme 11 The Zn(II) bis(formazanate) species across the three oxidation states.

Interestingly, Otten and co-workers also reported a structurally similar homoleptic formazanate Fe(II) complex, which adopted a low-spin iron(I) configuration upon one-electron reduction, rather than generating a ligand-centered radical (Scheme 12, top).59 Conversely, the one-electron reduction of a low-coordinate Fe(II) complex bearing a single formazanate ligand reported by Broere, Holland and co-workers (Scheme 12, bottom) resulted in a species that features a ground state with substantial multiconfigurational character with two dominant configurations. One represents a high-spin Fe(II) center that is antiferromagnetically coupled to a ligand-centered radical, whereas the other is best described as a high-spin Fe(I) complex.60 Despite having a significantly less negative reduction potential (0.8 V) in comparison to analogous b-diketiminate complexes, the anionic complex displayed similar reactivity in terms of halogen atom abstraction from alkyl iodides. On the other hand, the more accessible reduced state also enabled distinctive reactivity involving the reductive elimination of NdI bonds that was not seen for the b-diketiminate analog. Similar to the PDI systems described above, these iron formazanate systems present another example where it is challenging to identify whether redox events are ligand- or metal-centered, and also can be dependent on changes in geometry of the complex and presence of ancillary ligands. This poses a challenge for DFT calculations and can often only be captured using more advanced computational methods. Although formazanates are promising RALs, most work thus far has mainly focused on ligand synthesis, coordination chemistry and photochemistry of these systems.15 This has exposed interesting ligand rearrangements,61–63 and potential deleterious NdN scission side reactions that can occur upon two-electron reduction.59,61

Scheme 12 Top: metal-centered reduction in a bisformazanate Fe complex. Bottom: One-electron reduction of a low-coordinate Fe(II)-formazanate complex resulting in a complex with a multiconfigurational ground state, and subsequent reactivity.

In 2010, the Soper group reported that RALs on a Co(III) platform can be employed for CdC cross-coupling reactions.64 The anionic Co(III) bis(amidophenolate) complex depicted in Scheme 13 can react with alkyl halides such as CH2Cl2 or ethyl bromide to form a square pyramidal alkyl-Co(III) species. The two electrons necessary for the transformation are provided by the oxidation of both amidophenolate ligands to the respective iminosemiquinones. The authors proposed two possible pathways for the transformation, either through an electron transfer mechanism or a SN2 mechanism, although the data were in better agreement

Redox-Active Ligands in Organometallic Chemistry

429

Scheme 13 Electrophilic addition of an alkyl fragment to square planar Co(III) enabled by one-electron oxidation of two amidophenolate ligands.

with the latter. Exposing the square pyramidal ethyl-Co(III) species to organozinc reagents such as phenylzinc bromide and hexylzinc bromide led to the formation of ethylbenzene and n-octane respectively, concomitantly reforming the initial anionic species. Combining both transformations gives a net (non-catalytic) CdC cross-coupling reaction that formally does not involve metal-centered redox changes, as the electrons necessary for the transformations are provided by the RALs. The Soper group also reported that RALs can facilitate bimetallic O2 homolysis on pentacoordinate Re(V)-oxo complexes bearing two redox-active amidophenolate or catecholate ligands.65 The proposed mechanism (Scheme 14) starts by formation of a bound Z1-superoxo species with concomitant oxidation of one of the catecholate ligands to the semiquinone state, which avoids the formation of the an unfavorable Re(VI) center (A). The superoxide species is then intercepted by another equivalent of the starting Re(V) species, in which a catecholate ligand also undergoes one-electron oxidation, to form a dinuclear species with a peroxide bridge (B). The final Re(VII)-bis(oxo) species is then formed through homolytic cleavage of the bridging peroxide and reduction of the semiquinone ligands to the catecholate forms (C). An analogous complex bearing amidophenolato ligands showed the same reactivity and was proposed to proceed via a similar mechanism. The role of the RALs in this particular example was indicated by the observation that a structural analog bearing redox-inactive ligands did not show analogous reactivity toward O2.

Scheme 14 Dinuclear homolytic cleavage of O2 by a pentacoordinate Re(V) oxo bearing two RALs.

In 2005 Chaudhuri et al. reported that a square planar Cu(II) complex bearing two iminosemiquinonato ligands can oxidatively add Br2 to give a six-coordinate Cu(II) complex bearing two bromide ligands and two iminoquinone ligands (Scheme 15, top).66 This reactivity is similar to chemistry reported by Heyduk with Zr (Scheme 1), because the copper retains its original formal oxidation state as each RAL is oxidized by one electron. More recently, Desage-El Murr, Fensterbank and co-workers used the same

Scheme 15 Electron reservoir behavior of iminosemiquinonato ligands on Cu(II) in an oxidative addition (top) and synthesis of a cationic Cu-CF3 complex.

430

Redox-Active Ligands in Organometallic Chemistry

system to isolate the first example of a Cu(II) complex featuring a trifluoromethyl ligand.67 This was achieved through use of the Umemoto reagent (Scheme 15, bottom) resulting in the formation of a cationic Cu-CF3 complex. In this reaction the copper center retains its divalent oxidation state through one-electron oxidation of both iminosemiquinonato ligands to iminoquinone ligands. The resulting complex was shown to trifluoromethylate electrophiles, demonstrating the formal umpolung of the initially electrophilic CF3 group into a nucleophile. In later work the same group reported that the cationic Cu(II) complex bearing fully oxidized iminobenzoquinone ligands reacts as a masked Cu(III) reagent. Reactions of this complex with arylboronic acids result in N-arylation of one of the ligands.68 Mechanistic studies involving EPR spectroscopy and DFT calculations suggest a mechanism akin to a reductive elimination at the copper center. This demonstrates that redox-active ligands are capable of stabilizing/generating masked valence states, which can open up new avenues of reactivity. The same group demonstrated that the same square-planar Cu complex can also serve as a catalyst for controlled radical trifluoromethylation of various unsaturated substrates, including silyl enol ethers and heteroaromatics, using CF+3 sources.69 This reactivity is proposed to be enabled by the ability of only one RAL changing between changing between the iminosemiquinonato and iminobenzoquinone oxidation state. Interestingly, this single-electron transfer reactivity could also be translated to related Ni(II) complexes.70

1.13.2.3

Other metals

Similar to first-row transition metals, lanthanides and actinides typically undergo one-electron redox changes if anything. Hence chemists have also taken interest in the exploitation of RALs to enable multi-electron transformations on these metals with a particular focus on uranium. A variety of RALS, including a-diimines,71,72 2,20 -bipyridine,73–75 Z6-arenes,76,77 and iminoquinones78,79 have been used toward this end. In 2014, Bart and co-workers reported that redox-active PDI ligands bound to U(IV) can serve as an electron reservoir that can store up to three electrons for subsequent reactivity.80 The group also showed that related complexes featuring the PDI ligand in the highly reduced tris-anionic state can oxidatively add dihalides, selenylchlorides and dichalcogenides (Scheme 16).81 In all cases the PDI ligand was oxidized by two electrons to the corresponding monoanion, thereby keeping the uranium center in the same oxidation state.

Scheme 16 Oxidative addition on U(IV) bearing a triply reduced PDI ligand.

In 2015, the same group reported related complexes for the scission of U-O bonds,82 which is a challenging transformation of relevance in the processing and remediation of uranium in the environment. More recently, Bart and co-workers reported breaking UdO bonds in uranyl compounds by a reaction with pivaloyl chloride (Scheme 17) or B-chlorocatecholborane.83 Detailed spectroscopic studies and isotopic labeling showed that the O atoms in the respective dicatecholato-bisborylether and pivaloyl anhydride originate from the uranyl fragment. Interestingly, these reactions result in the reduction of U(VI) to U(IV) concomitant with one-electron oxidation of each iminosemiquinonato ligand.

Scheme 17 RAL-facilitated UdO bond scission using pivaloyl chloride.

Redox-Active Ligands in Organometallic Chemistry

1.13.3

431

Modification of the Lewis acid-base properties through ligand-centered redox changes

The ability of RALs to reversible accept or donate electrons is an appealing strategy to generate additional accessible redox states. As showcased in the previously section, RALs can undergo redox changes independently of the metal that they are bound to. The associated changes in ligand binding mode also directly affect the Lewis acidity or basicity of the bound metal center, thereby affecting the coordination behavior of a bound substrate or hemilabile donor. An early example of this was reported by the group of Rauchfuss who reported the synthesis of a square planar Pt(II) complex bearing a redox-active amidophenolate and cyclooctadiene ligand. They found that the oxidation state of the RAL strongly affects the degree to which the cyclooctadiene is susceptible toward nucleophilic attack.84 The neutral complex was unreactive toward nucleophilic attack by alkoxides. However, a ligand-centered one-electron oxidation activated the bound diene toward nucleophilic attack by sodium methoxide or ethoxide (Scheme 18). Interestingly, the reaction was highly stereospecific toward CdO bond formation at the vinylic carbon trans to the O atom of the amidophenolate ligand. The reaction is reversible through the addition of HPF6, yielding the cationic Pt(II) complex and methanol or ethanol. The ability to switch a catalyst on or off through ligand-centered oxidation could provide opportunities for orthogonal reactivity patterns toward different substrates in cascade reactions.

Scheme 18 Redox-switchable nucleophilic attack on a COD-Pt complex.

The same group also reported similar redox-switchable behavior with a coordinatively unsaturated, 16 valence electron Cp∗Ir complex bearing a redox-active amidophenolate ligand (Scheme 19).85 One-electron oxidation of the RAL to the iminosemiquinone state (A) increases the Lewis acidity enabling reversible binding of H2 (B). In the presence of non-coordinating base and excess oxidant this enabled catalytic H2 oxidation (C). In the absence of a base, the complex decomposed to the dimeric [Cp∗Ir(m2-H)3IrCp∗]+ species and free protonated aminophenol ligand. Follow-up studies of the kinetics indicated amongst others a counterion dependence and demonstrated that the Cp∗Rh analog displayed the same behavior.86 The Sarkar group also reported similar redox-switchable KOtBu-assisted H2 activation with amidophenolate-Pt(II)pap (pap ¼ phenylazopyridine) system.87

Scheme 19 Redox-switchable H2 oxidation on a Cp∗Ir complex involving ligand-centered redox changes

432

Redox-Active Ligands in Organometallic Chemistry

Another way that reactivity could potentially be “switched on/off” is through ligand-centered redox changes that affect the coordination of hemilabile donors. Kaim and coworkers demonstrated this behavior with coordinatively unsaturated, 16 valence electron (Ru and Os) complexes bearing amidophenolate-derived ligands.88 A series of amidophenolate ligands featuring different hemilable donor substituents were investigated (Scheme 20), but in no neutral compound was coordination of the hemilabile donor atom observed. However, one-electron oxidation of the RAL to the iminosemiquinone oxidation state resulted in coordination of the thio- and selenoether moieties, which was shown by X-ray crystallography. Although no single crystals of the ether analog were obtained, the authors argue that the ether behaves in a similar fashion as the thio- and selenoether analogs, based on spectroelectrochemical experiments. Similar behavior was found for the aforementioned a low-coordinate Fe formazanate complex (Scheme 12), which displays strong THF binding that is lost upon one-electron reduction of the RAL.59

Scheme 20 Ligand-centered redox changes affect ligand hemilability for a Ru(p-cymene) complex.

1.13.4

Redox-active ligand-to-substrate single-electron transfer

The utilization of metal-bound reactive carbene and nitrene-centered radicals has been demonstrated as a powerful tool to enable CdC and CdN bond formation in controlled radical catalysis.89 These reactive substrate-centered radicals are typically formed by a reaction of a carbene or nitrene precursor with a redox-active metal center (e.g. Fe(II) or Co(I)), concomitant with single-electron transfer from the metal to the carbene/nitrene. Interestingly, RALs can also serve as the electron donor to generate such reactive substrate-centered radicals. Many of the examples above employ redox-active ligands to facilitate two-electron reactivity with first-row transition metals. In contrast, van der Vlugt and co-workers showed that RALs can also enable single-electron reactivity on a metal well known for its two-electron transformations: Pd. A reaction of a redox-active aminophenol-derived NNO ligand with PdCl2 provided access to a square planar Pd(II) complex featuring a mono-anionic iminosemiquinonato radical ligand (Scheme 21), as shown by various experimental and computational analyses.90 One-electron reduction using CoCp2 provided a diamagnetic complex, which was found to convert 1-(4-azido-butyl)benzene into the corresponding pyrrolidine analogs. Mechanistic investigations involving DFT calculations, isotopic labelling and trapping experiments suggested single-electron transfer from the RAL to the nitrene generates an intermediate featuring a reactive nitrene radical. A subsequent intramolecular H atom abstraction and radical rebound produces the pyrrolidine product, which reacts with Boc-anhydride (Boc2O) to form the Boc-protected pyrrolidine product. The proposed mechanism for this transformation is similar to mechanisms proposed for the Fe(II) or Co(I) catalyst mentioned above, but the distinct difference is that no metal-centered redox changes are involved.

Scheme 21 Radical CdH amination of organic azides to pyrrolidines mediated by ligand-to-substrate single-electron transfer on a Pd(II) center.

Although conceptually interesting, the turnover number (TON) of the NNOPd system was poor.91 In contrast, the analogous NNOFeCl2 complex proved a more efficient catalyst for CdH amination of unactivated organic azides to saturated N-heterocycles, with TON’s up to 620.92 The same group found that similar coordination chemistry was observed when replacing the pyridine donor in the NNO ligand by a arylphosphine fragment (Scheme 22). However, this significantly lowered the reduction potential, which is in agreement with a ligand-centered process.93 A reaction of the reduced complex with disulfides, in the presence of

Redox-Active Ligands in Organometallic Chemistry

433

Scheme 22 Homolytic cleavage of PhSSPh enabled by ligand-centered oxidation, and formation of the thiolate-bridged dipalladium(II) complex.

TlPF6 to abstract the chloride, resulted in homolytic cleavage of the disulfide bond yielding a dinuclear complex with ligand-based mixed valence. Crossover experiments suggested the generation of a thiyl radical upon reaction with the Pd complex, which undergoes a ligand-centered one-electron oxidation. This example of homolytic cleavage of a disulfide facilitated by a RAL differs from the more common two-electron oxidative addition of disulfides on a metal center and demonstrates how redox-active ligands can give access to new modes of reactivity. Another example of ligand-to-substrate single-electron transfer to generate reactive nitrene radicals was reported by Desage-El Murr and coworkers in 2018.94 To this end they utilized a neutral Cu(II) complex bearing two iminosemiquinonato radical anion ligands (Scheme 23).2 This complex has an overall S ¼ ½ ground state due to strong antiferrogmagnetic coupling of one of the

Scheme 23 Catalytic aziridination of 4-chlorostyrene by a Cu(II) bis-iminosemiquinonato complex.

434

Redox-Active Ligands in Organometallic Chemistry

ligand radicals with the copper-centered radical. Treatment of this Cu(II) complex with iminoiodinanes, which are common nitrene precursors, results in the generation of a reactive nitrene radical concomitant with oxidation of one of the RALs to the iminobenzoquinone oxidation state. This gives rise to a quartet ground state, but the RALs and Cu are proposed to generate an intermediate with a low-lying doublet state, thereby enabling this otherwise formally spin-forbidden transformation. This system was shown to be a competent catalyst for aziridination of various types of alkenes. Notably, mechanistic studies using EPR spectroscopy to monitor the reaction of the nitrene radical complex with 4-chlorostyrene enabled the observation of the corresponding insertion product featuring a benzylic radical (Scheme 23, bottom). A final radical rebound results in aziridine formation and regeneration of the active catalyst. A final example of ligand-to-substrate single electron transfer involves the TAML (tetra-amido macrocyclic ligand) scaffold, which has found widespread use in oxidation catalysis. The redox-activity of the TAML platform on Co(III) has, until recently,95 been under debate. In a comprehensive experimental and computational study, de Bruin and co-workers demonstrated that one-electron oxidation of the anionic complex depicted in Scheme 24 is indeed ligand-centered. Although their DFT calculations reproduced earlier findings, they found that the oxidation product had substantial multireference character, which is often poorly described using single-reference DFT calculations. The authors found that both complexes reacted with nitrene precursors to give nitrene radical complexes. However, a remarkable finding is that the ligand oxidation state determines whether a mono- or bis-nitrene radical complex is formed associated with a single or double ligand-to-substrate single-electron transfer, respectively. Both nitrene radical complexes were found to be effective catalysts for diastereoselective aziridination of styrene derivatives, but detailed mechanistic studies showed that they do not react via the expected radical addition and radical rebound mechanism.96 Instead their findings suggest a new type of mechanism wherein CdN bond formation happens through an unusual electronically asynchronous transition state. This transition state more closely resembles initial oxidation of the styrene by the CoTAML fragment, followed by a nucleophilic attack of the nitrene lone pair on a styrene radical cation (Scheme 25).

Scheme 24 Formation of bis- and mono- nitrene Co(III)TAML complexes.

Scheme 25 Electronically asynchronous transition states during catalytic aziridination of styrene by a Co(III)TAML-nitrene complex.

Redox-Active Ligands in Organometallic Chemistry

1.13.5

435

Cooperative ligand-centered reactivity

The previous sections feature RALs that are supporting ligands, and are thus designed to prevent ligand-centered chemical transformations beyond redox changes. Reactions at the radical-bearing ligand are usually inhibited using sterically demanding moieties on the RAL and/or a large p-system for delocalization of unpaired electrons. As described in Section 1.13.4, RALs can also generate substrate-centered radicals for inter- or intramolecular reactions to form an product that dissociates from the metal complex. In this section, selected examples will be described where the supporting RAL plays a more active role by employing ligand-centered radicals in reversible bond making and breaking processes. The natural metalloenzyme galactose oxidase (GOase), which performs the aerobic oxidation of primary alcohols to aldehydes and H2O2 with high turnover frequencies, is one of the most well-studied enzymes wherein a RAL plays such an active role. The active site of GOase contains a Cu(II) center bound to two histidine residues, a thioether-linked tyrosine residue in the equatorial plane and a tyrosine in the axial position (Scheme 26).97 The latter operates as a base throughout the proposed catalytic cycle, but it is the thioether-linked tyrosine that fulfills the role of a redox-active actor ligand. Initially the inactive resting state of the metalloenzyme is activated by single electron oxidation, which results in a ligand-centered radical on the thioether-linked tyrosine instead of a Cu(III) center (A). Subsequently the axial tyrosine deprotonates the alcohol substrate which displaces the water molecule (B). With the formed alkoxide bound to the Cu(II) center, proton coupled electron transfer (PCET) from the alkoxide b-hydrogens to the tyrosine radical takes place, resulting in a copper bound ketyl radical (C). The oxidation of the ketyl to an aldehyde and subsequent dissociation from the metal center results in reduction of the copper center to Cu(I) (D). This reduced and coordinatively unsaturated state of the active site can then react with dioxygen to give a Cu(II) bound superoxide (E). PCET from the thioether-linked tyrosine to the superoxide regenerates the tyrosine-centered radical and yields a copper bound hydroperoxide (F). Finally, a reaction with water yields hydrogen peroxide and the initial active species (G).

Scheme 26 The proposed mechanism for the catalytic aerobic oxidation of primary alcohols to aldehydes and H2O2 by GOase.

Inspired by the GOase enzyme, Wieghardt and co-workers developed synthetic Cu and Zn analogs of the active site of the enzyme and demonstrated the catalytic aerobic oxidation of primary alcohols.98 This was achieved by using a tetradentate

436

Redox-Active Ligands in Organometallic Chemistry

bis(aminophenol) ligand, serving as the structural model of the active site. The ligand is readily synthesized through condensation of 3,5-di-tert-butylcatechol and o-phenylenediamine in a 2:1 ratio. The divalent copper and zinc complexes of this tetradentate ligand were found to be active catalysts for aerobic oxidation of primary alcohols (a feat other GOase structural models were unable to achieve up until that point).99,100 Based on the observed kinetics of the stoichiometric and catalytic transformations, including the analysis of kinetic isotope effects (KIEs), the authors proposed the mechanism depicted in Scheme 27, which is very similar to the proposed mechanism of GOase. First, the alcohol is deprotonated by one of the phenolate oxygen atoms and the resulting alkoxide then coordinates to the divalent metals (A). Subsequently, the rate-determining H-atom abstraction takes place by the phenoxyl radical, generating the ketyl radical (B). This is followed by single-electron transfer from the ligand backbone to the ketyl radical, concomitant with dissociation of the aldehyde (C). Finally, a reaction with O2 regenerates the starting complex and gives and releases H2O2 (D). Although the same mechanism was proposed for both systems, the turnover frequency (TOF) of the Cu-based catalyst ( 0.03 s-1) far exceeds that of the Zn-analog (0.002 s−1). This difference was also reflected in the observed rate constant for the intramolecular H-atom abstraction (step B), which was faster by a factor of 10 for the Cu-based system.

Scheme 27 The proposed mechanism for the catalytic aerobic oxidation of primary alcohols by the galactose oxidase structural models.

In related work, Chaudhuri, Wieghardt and coworkers developed a sulfur-bridged bisphenolate-based dicopper(II) complex that is capable of oxidizing alcohols to the corresponding aldehydes using O2 as the oxidant.32 Interestingly, when using secondary alcohols, the formation of the 1,2-glycols was observed (Scheme 28). Both reactions were found to obey rate laws that differ in the reaction order of alcohol concentration. Similar to the previous examples, KIEs larger than 8 were observed, indicating that hydrogen abstraction is the rate determining step. Based on experimental observations, the authors proposed a catalytic cycle wherein the RAL plays a similar role as the thioether-linked tyrosine in GOase, yet without involving Cu-centered redox changes. First the dicopper species reduces O2, forming H2O2 and a dinuclear copper species with two ligand-centered radicals (A). Subsequently, the m-alkoxides deprotonate a secondary alcohol to give an alkoxide on each of the copper ions (B). In the next step, the two phenoxyl radicals abstract the b-hydrogens from the alkoxides – though there is no experimental data supporting that this happens simultaneously – (C) leading to the formation of two copper-bound ketyl moieties. The latter are proposed to couple to release the corresponding 1,2-glycol and reform the initial dicopper(II) complex (D). A noteworthy observation is that methanol is not a suitable substrate, and the authors attribute this to the need for an alcohol substrate with relatively weak CdH bonds to enable CdH abstraction from the coordinated alkoxide.

Redox-Active Ligands in Organometallic Chemistry

437

Scheme 28 Catalytic 1,2-glycol formation using a sulfur-bridged bisphenolate-based dicopper(II) complex.

Although the previous examples mimic specific features of the operating mechanism that GOase uses to oxidize alcohols, their catalytic activity is low. Similarly inspired by GOase, the Grützmacher group developed a bis(dibenzotropylamino)-based Ir system capable of catalytically oxidizing primary alcohols to aldehydes. This system was found to display very high activity with TOFs > 40 s-1, albeit using a scarce metal and benzoquinone as the oxidant.101 The catalytic cycle (Scheme 29) proposed by the authors starts from the neutral aminoamidate complex, which can be deprotonated by potassium tert-butoxide to give the anionic bis(amidate) complex (A). One-electron oxidation of the anionic complex by benzoquinone leads to the formation of a neutral, radical complex (B) that can react with the primary alkoxides, forming an anionic Ir-alkoxide complex with an aminyl radical (C). The aminyl radical can then abstract a hydrogen atom from the a-carbon of the alkoxide, giving the Ir-ketyl species (D). This species can then be oxidized by the semiquinonate to liberate the aldehyde and regenerate the starting aminoamidate species (E). A final example of how redox-active actor ligands can be employed was initially reported by Wang and Stiefel in 2001, who exploited ligand-centered redox for the separation and purification of olefins. They described an electrochemical approach that employed a Ni(II) bis-dithiolene complex to separate light olefins from a multi-component stream (MCS) containing various other gases (Scheme 30).102 The authors propose that upon electrochemical oxidation of the anionic Ni(II) bis-dithiolene complex to the neutral analog, it can form an adduct with the olefin bridging two S atoms from the dithiolene ligands. Subsequent electrochemical reduction of this adduct first forms the anionic analog, which can then release the olefin and regenerate the initial anionic Ni(II) bis-dithiolene complex. Mechanistic studies from Fekl demonstrated that in the absence of the anionic bis-dithiolene complex, the reaction between the neutral complex and lighter olefins leads to the formation of dihydrodithiins and metal decomposition

438

Redox-Active Ligands in Organometallic Chemistry

Scheme 29 The proposed mechanism for the catalytic oxidation of primary alcohols to aldehydes by a bis(dibenzotropylamino)-based Ir complex.

Scheme 30 Separation of ethylene from a MCS by a Ni(II) bis-dithiolene complex.

products, whereas in the presence of the anionic complex the neutral analog reacts with olefins to predominantly form the interligand dithiolene-olefin adducts.103 Combined experimental and theoretical work were indicative of the anionic complex catalyzing the formation of the interligand adduct by formation of a bimetallic intermediate, which suppresses the undesired reaction of the neutral complex with olefins to form dihydrodithiins alongside various nickel-containing decomposition products.104

Redox-Active Ligands in Organometallic Chemistry

1.13.6

439

Conclusion

What started as a controversy over how to describe the electronic structure of a series of nickel complexes has developed into an exciting field in coordination and organometallic chemistry and catalysis. The examples described in this chapter clearly show that redox-active ligands can bring about new reactivity in various ways. These ligands can have a “spectator” role where they mediate metal-centered reactivity by changing the Lewis acidity as a function of ligand oxidation state (Section 1.13.3), avoid unfavorable metal oxidation states, or by serving as an electron reservoir that provides/accepts electrons needed for metal-centered bond making/breaking events. In Section 1.13.2, several examples are described where the latter enables the characteristic two-electron transformations of noble metals with more earth-abundant first row transition metals that typically undergo one-electron redox changes, or with metals that have no formal d-electrons at all. A relatively recent concept exploits the possibility to transfer a single electron from a redox-active spectator ligand to a redox-active substrate (Section 1.13.4). This approach, which does not rely on metal-centered redox changes, has been demonstrated to enable radical-type transformations using metals that typically do not undergo single-electron changes. In addition, it can be used to change the mechanism of key chemical transformations. This provides the opportunity of mediating unprecedented chemical transformations that benefit from the different coordination behavior on metals in specific oxidation states that typically do not enable substrate-centered radical-type reactivity. Finally, chemists have also developed systems wherein RALs can behave as “actor” ligands. This approach is inspired by naturally occurring metalloenzymes, wherein stabilized ligand-centered radicals are exploited to enable the reversible formation of covalent bonds with a substrate. In this way it is possible to achieve different chemoselectivity, to accelerate reaction steps involving the transfer of both protons and electrons or to enable new, substrate-specific binding modes (Section 1.13.5). The presence of additional low-energy molecular orbitals, which enables RALs to expand on a metal’s common reactivity, can also lead to a series of electronic states that are densely spaced in their relative energies. Such a high density of states105 can result in multiconfigurational ground states and also facilitate two-state reactivity wherein multiple spin surfaces participate in reaction pathways.106 In Sections 1.13.2 and 1.13.4 several cases are described wherein DFT calculations fail to accurately describe the multiconfigurational electronic structures of first-row transition metals bearing RALs. Moreover, an example is described where distinct reactivity was long ascribed to the RAL acting as an electron reservoir, but recent studies revealed substrate dependent two-state reactivity. Unfortunately, rational design of such systems can be hampered by multiconfigurational ground states, which are challenging to identify and require computationally expensive multireference methods to model. On the other hand, in recent years the use of multireference computational methods for studying metal complexes featuring RALs is becoming more common (for example see Refs. 49, 60, 95, 96), and the widespread usage of such methods should be a real gamechanger. In summary, RALs can push metals out of their comfort zone into uncharted territory, yet they can do so by leaving them in their comfortable oxidation states. They are one of the latest tools in the organometallic toolbox and further exploration of their potential will undoubtedly lead to exciting new developments in organometallic chemistry and catalysis.

Acknowledgment We thank Dr. Klaas van Leest for insightful discussions that aided in the preparation of this chapter. Roel Bienenmann and Cody van Beek are acknowledged for their thorough proofreading.

References 1. For example see (a) DiMucci, I.D.; Lukens, T. T.; Chatterjee, S.; Carsch, K. M.; Titus, C. J.; Lee, S.; Nordlund, D.; Betley, T. A.; Macmillan, S. N.; Lancaster, J. Am. Chem. Soc. 2019, 141, 18508–18520;(b) Steen, J. S.; Knizia, G.; Klein, J. E. M. N. Angew. Chem. 2019, 131, 13267–13273; (c) Ye, S.; Geng, C.-Y.; Shaik, S.; Neese, F. Phys. Chem. Chem. Phys. 2013, 15, 8017–8030 2. Chaudhuri, P.; Verani, C. N.; Bill, E.; Bothe, E.; Weyhermüller, T.; Wieghardt, K. J. Am. Chem. Soc. 2001, 123, 2213–2223. 3. (a) Schrauzer, G. N.; Mayweg, V. J. Am. Chem. Soc. 1962, 84 (16), 3221; (b) Gray, H. B.; Williams, R.; Bernal, I.; Billig, E. J. Am. Chem. Soc. 1962, 84, 3596–3597; (c) Gray, H. B.; Billig, E. J. Am. Chem. Soc. 1963, 85, 2019–2020; (d) Davison, A.; Edelstein, N.; Holm, R. H.; Maki, A. H. Inorg. Chem. 1963, 2, 1227–1232; (e) Locke, J.; McCleverty, J. A.; Wharton, E. J.; Winscom, C. J. Chem. Commun. 1966, 19, 677–678. 4. (a) Eisenberg, R.; Gray, H. B. Inorg. Chem. 2011, 50, 9741–9751; (b) Hine, F. J.; Taylor, A. J.; Garner, C. D. Coord. Chem. Rev. 2010, 254, 1570–1579. 5. Jørgensen, C. K. Inorg. Chim. Acta Rev. 1968, 2 (C), 65–88. 6. Kaim, W. Inorg. Chem. 2011, 50, 9752–9765. 7. Chirik, P. J. Inorg. Chem. 2011, 50, 9737–9740. 8. (a) Adhikari, D.; Mossin, S.; Basuli, F.; Huffman, J. C.; Szilagyi, R. K.; Meyer, K.; Mindiola, D. J. J. Am. Chem. Soc. 2008, 130, 3676–3682; (b) Harkins, S. B.; Mankad, N. P.; Miller, A. J. M.; Szilagyi, R. K.; Peters, J. C. J. Am. Chem. Soc. 2008, 130, 3478–3485; (c) Radosevich, A. T.; Melnick, J. G.; Stoian, S. A.; Bacciu, D.; Chen, C. H.; Foxman, B. M.; Ozerov, O. V.; Nocera, D. G. Inorg. Chem. 2009, 48, 9214–9221; (d) Vreeken, V.; Siegler, M. A.; de Bruin, B.; Reek, J. N. H.; Lutz, M.; van der Vlugt, J. I. Angew. Chem. Int. Ed. 2015, 54, 7055. Angew. Chem. 2015, 127, 7161; (e) Wang, D.; Lindeman, S. V.; Fiedler, A. T. Inorg. Chem. 2015, 54, 8744–8754; (f ) Vreeken, V.; Broere, D. L. J.; Jans, A. C. H.; Lankelma, M.; Reek, J. N. H.; Siegler, M. A.; van der Vlugt, J. I. Angew. Chem. Int. Ed. 2016, 55, 10042–11046. Angew. Chem. 2016, 128, 10196–10200; (g) Vreeken, V.; Baij, L.; de Bruin, B.; Siegler, M. A.; van der Vlugt, J. I. Dalton Trans. 2017, 46, 7145–7149; (h) Vreeken, V.; Siegler, M. A.; van der Vlugt, J. I. Chem. Eur. J. 2017, 23, 5585–5594; (i) Leconte, N.; Moutet, J.; Herasymchuk, K.; Clarke, R. M.; Philouze, C.; Luneau, D.; Storr, T.; Thomas, F. Chem. Commun. 2017, 53, 2764–2767. 9. (a) Heins, S. P.; Wolczanski, P. T.; Cundari, T. R.; MacMillan, S. N. Chem. Sci. 2017, 8, 3410–3418; (b) Villa, M.; Miesel, D.; Hildebrandt, A.; Ragaini, F.; Schaarschmidt, D.; Jacobi von Wangelin, A. ChemCatChem. 2017, 9, 3203–3209.

440

Redox-Active Ligands in Organometallic Chemistry

10. (a) Caulton, K. G. Eur. J. Inorg. Chem. 2012, 435–443; (b) Budzelaar, P. H. M. Eur. J. Inorg. Chem. 2012, 3, 530–534; (c) Sieh, D.; Schlimm, M.; Andernach, L.; Angersbach, F.; Neckel, S.; Schöffel, J.; Šušnjar, N.; Burger, P. Eur. J. Inorg. Chem. 2012, 3, 444–462; (d) Myers, T. W.; Berben, L. A. Inorg. Chem. 2012, 51, 1480–1488; (e) Chirik, P. J. Angew. Chem. Int. Ed. 2017, 56, 5170–5181; (f ) Williams, V. A.; Hulley, E. B.; Wolczanski, P. T.; Lancaster, K. L.; Lobkovsky, E. B. Chem. Sci. 2013, 4, 3636–3648; (g) Bowman, A. C.; Tondreau, A. M.; Lobkovsky, E.; Margulieux, G. W.; Chirik, P. J. Inorg. Chem. 2018, 57, 9634–9643; (h) Wang, Q.; Zhang, S.; Cui, P.; Weberg, A. B.; Thierer, L. M.; Manor, B. C.; Gau, M. R.; Carroll, P. J.; Tomson, N. C. Inorg. Chem. 2020, 59, 4200–4214; (i) Zhang, S.; Cui, P.; Liu, T.; Wang, Q.; Longo, T. J.; Thierer, L. M.; Manor, B. C.; Gau, M. R.; Carroll, P. J.; Papaefthymiou, G. C.; Tomson, N. C. Angew. Chem. Int. Ed. 2020, 59, 15215–15219. 11. (a) Pierpont, C. G. Coord. Chem. Rev. 2001, 219–221, 415–433; (b) Pierpont, C. G. Coord. Chem. Rev. 2001, 216–217, 99–125; (c) Pierpont, C. G. Inorg. Chem. 2011, 50, 9766–9772; (d) Mederos, A.; Domínguez, S.; Hernández-Molina, R.; Sanchiz, J.; Brito, F. Coord. Chem. Rev. 1999, 193–195, 857–911. 12. (a) Ciccione, J.; Leconte, N.; Luneau, D.; Philouze, C.; Thomas, F. Inorg. Chem. 2016, 55, 649–665; (b) van der Meer, M.; Manck, S.; Sobottka, S.; Plebst, S.; Sarkar, B. Organometallics 2015, 34, 5393–5400; (c) Mederos, A.; Domínguez, S.; Hernández-Molina, R.; Sanchiz, J.; Brito, F. Coord. Chem. Rev. 1999, 913–939; (d) Nguyen, A. I.; Blackmore, K. J.; Carter, S. M.; Zarkesh, R. A.; Heyduk, A. F. J. Am. Chem. Soc. 2009, 131, 3307–3316; (e) Munhá, R. F.; Zarkesh, R. A.; Heyduk, A. F. Inorg. Chem. 2013, 52, 11244–11255; (f ) Ghosh, S.; Baik, M.-H. Chem. Eur. J. 2015, 21, 1780–1789. 13. (a) Blackmore, K. J.; Ziller, J. W.; Heyduk, A. F. Inorg. Chem. 2005, 44, 5559–5561; (b) Zarkesh, R. A.; Heyduk, A. F. Organometallics 2011, 30, 4890–4898; (c) Hananouchi, S.; Krull, B. T.; Ziller, J. W.; Furche, F.; Heyduk, A. F. Dalton Trans. 2014, 43, 17991–18000. 14. (a) Gilroy, J. B.; Koivisto, B. D.; McDonald, R.; Ferguson, M. J.; Hicks, R. G. J. Mater. Chem. 2006, 16, 2618–2624; (b) Gilroy, J. B.; McKinnon, S. D. J.; Koivisto, B. D.; Hicks, R. G. Org. Lett. 2007, 9, 4837–4840; (c) McKinnon, S. D. J.; Patrick, B. O.; Lever, A. B. P.; Hicks, R. G. Chem. Commun. 2010, 46, 773–775; (d) Sanz, C. A.; Ferguson, M. J.; McDonald, R.; Patrick, B. O.; Hicks, R. G. Chem. Commun. 2014, 50, 11676–11678; (e) Johnston, C. W.; Schwantje, T. R.; Ferguson, M. J.; McDonald, R.; Hicks, R. G. Chem. Commun. 2014, 50, 12542–12544; (f ) Sanz, C. A.; McKay, Z. R.; MacLean, S. W. C.; Patrick, B. O.; Hicks, R. G. Dalton Trans. 2017, 46, 12636–12644. 15. Gilroy, J. B.; Otten, E. Chem. Soc. Rev. 2020, 49, 85–113. 16. Nippe, M.; Khnayzer, R. S.; Panetier, J. A.; Zee, D. Z.; Olaiya, B. S.; Head-Gordon, M.; Chang, C. J.; Castellano, F. N.; Long, J. R. Chem. Sci. 2013, 4, 3934–3945. 17. (a) Lichtenberg, C.; Krommenacher, I. Chem. Commun. 2016, 52, 10044–10047; (b) Hanft, A.; Lichtenberg, C. Dalton Trans. 2018, 47, 10578–10589; (c) Hanft, A.; Lichtenberg, C. Eur. J. Inorg. Chem. 2018, 3361–3373. 18. (a) Doslik, N.; Sixt, T.; Kaim, W. Angew. Chem. Int. Ed. 1998, 37, 2403–2404. Angew. Chem. 1998, 110, 2521–2522; (b) Dogan, A.; Sarkar, B.; Klein, A.; Lissner, F.; Schleid, T.; Fiedler, J.; Záliš, S.; Jain, V. K.; Kaim, W. Inorg. Chem. 2004, 43, 5973–5980; (c) Patra, S.; Sarkar, B.; Maji, S.; Fiedler, J.; Urbanos, F. A.; Jimenez-Aparicio, R.; Kaim, W.; Lahiri, G. K. Chem. Eur. J. 2006, 12, 489–498; (d) Paul, N. D.; Samanta, S.; Goswami, S. Inorg. Chem. 2010, 49, 2649–2655; (e) Joy, S.; Kr-mer, T.; Paul, N. D.; Banerjee, P.; McGrady, J. E.; Goswami, S. Inorg. Chem. 2011, 50, 9993–10004; (f ) Jana, R.; Lissner, F.; Schwederski, B.; Fiedler, J.; Kaim, W. Organometallics 2013, 32, 5879–5886; (g) Ghosh, P.; Samanta, S.; Roy, S. K.; Demeshko, S.; Meyer, F.; Goswami, S. Inorg. Chem. 2014, 53, 4678–4686; (h) Sengupta, D.; Ghosh, P.; Chatterjee, T.; Datta, H.; Paul, N. D.; Goswami, S. Inorg. Chem. 2014, 53, 12002–12013; (i) Sengupta, D.; Chowdhury, N. S.; Samanta, S.; Ghosh, P.; Seth, S. K.; Demeshko, S.; Meyer, F.; Goswami, S. Inorg. Chem. 2015, 54, 11465–11476; (j) Sinha, S.; Das, S.; Sikari, R.; Parua, S.; Brandal, P.; Demeshko, S.; Meyer, F.; Paul, N. D. Inorg. Chem. 2017, 56, 14084–14100; (k) Roy, S.; Pramanik, S.; Patra, S. C.; Adhikari, B.; Mondal, A.; Ganguly, S.; Pramanik, K. Inorg. Chem. 2017, 56, 12764–12774. 19. (a) Bezpalko, M. W.; Foxman, B. M.; Thomas, C. M. Inorg. Chem. 2013, 52, 12329–12331; (b) Harris, C. F.; Bayless, M. B.; van Leest, N. P.; Bruch, Q. J.; Livesay, B. N.; Bacsa, J.; Hardcastle, K. I.; Shores, M. P.; de Bruin, B.; Soper, J. D. Inorg. Chem. 2017, 56, 12421–12435; (c) Mahoney, J. K.; Martin, D.; Moore, C. E.; Rheingold, A. L.; Bertrand, G. J. Am. Chem. Soc. 2013, 135, 18766–18769; (d) Munz, D.; Chu, J.; Melaimi, M.; Bertrand, G. Angew. Chem. Int. Ed. 2016, 55, 12886–12890; (e) Eberle, B.; Kaifer, E.; Himmel, H.-J. Angew. Chem. Int. Ed. 2017, 56, 3360–3363. Angew. Chem. 2017, 129, 3408–3412; (f ) Taylor, J. W.; McSkimming, A.; Guzman, C. F.; Harman, W. H. J. Am. Chem. Soc. 2017, 139, 11032–11035; (g) Rosen, E. L.; Varnado, C. D., Jr.; Tennyson, A. G.; Khramov, D. M.; Kamplain, J. W.; Sung, D. H.; Creswell, P. T.; Lynch, V. M.; Bielawski, C. W. Organometallics. 2009, 28, 6695–6706; (h) Weinberg, D. R.; Hazari, N.; Labinger, J. A.; Bercaw, J. E. Organometallics 2010, 29, 89–100; (i) Romain, C.; Choua, S.; Colllin, J.-P.; Heinrich, M.; Bailly, C.; Karmazin-Brelot, L.; Bellemin-Laponnaz, S.; Dagorne, S. Inorg. Chem. 2014, 53, 7371–7376; (j) Hettmanczyk, L.; Suntrup, L.; Klenk, S.; Hoyer, C.; Sarkar, B. Chem. Eur. J. 2017, 23, 576–585; (k) Schrempp, D. F.; Schneider, E.; Kaifer, E.; Wadepohl, H.; Himmel, H.-J. Chem. Eur. J. 2017, 23, 11636–11648; (l) Ruamps, M.; Bastin, S.; Rechignat, L.; Saquet, A.; Valyaev, D. A.; Mouesca, J.-M.; Lugan, N.; Maurel, V.; Cesar, V. Chem. Commun. 2018, 54, 7653–7656; (m) Jongbloed, L. S.; Vogt, N.; Sandleben, A.; de Bruin, B.; Klein, A.; van der Vlugt, J. I. Eur. J. Inorg. Chem. 2018, 20-21, 2408–2418. 20. Fedushkin, I. L.; Dodonov, V. A.; Skatova, A. A.; Sokolov, V. G.; Piskunov, A. V.; Fukin, G. K. Chem. Eur. J. 2018, 24, 1877–1889. 21. Johansson Seechurn, C. C. C.; Kitching, M. O.; Colacot, T. J.; Snieckus, V. Angew. Chem. Int. Ed. 2012, 51, 5062–5085. 22. Stradiotto, M., Lundgren, R. J., Eds.; In Ligand Design in Metal Chemistry: Reactivity and Catalysis; John Wiley & Sons, 2016. 23. Chirik, P. J.; Wieghardt, K. Science 2010, 327, 794. 24. Kaim, W.; Schwederski, W. Coord. Chem. Rev. 2010, 254, 1580–1588. 25. Praneeth, V. K. K.; Ringenberg, M. R.; Ward, T. R. Angew. Chem. Int. Ed. 2012, 51, 10228–10234. 26. Luca, O. R.; Crabtree, R. H. Chem. Soc. Rev. 2012, 42, 1440–1459. 27. For a review specifically on this topic see: van der Vlugt, J. I. Chem. Eur. J. 2019, 25, 2651–2662. 28. Kuijpers, P. F.; van der Vlugt, J. I.; Schneider, S.; de Bruin, B. Chem. Eur. J. 2017, 23, 13819–13829. 29. Brown, S. N. Inorg. Chem. 2012, 51, 1251–1260. 30. Haneline, M. R.; Heyduk, A. F. J. Am. Chem. Soc. 2006, 128, 8410–8411. 31. Ketterer, N. A.; Fan, H.; Blackmore, K. J.; Yang, X.; Ziller, J. W.; Baik, M.; Heyduk, A. F. J. Am. Chem. Soc. 2008, 130, 4364–4374. 32. Chaudhuri, P.; Hess, M.; Flörke, U.; Wieghardt, K. Angew. Chem. Int. Ed. 1998, 37, 2217–2220. 33. Blackmore, K. J.; Lal, N.; Ziller, J. W.; Heyduk, A. F. J. Am. Chem. Soc. 2008, 130, 2728–2729. 34. Heyduk, A. F.; Zarkesh, R. A.; Nguyen, A. I. Inorg. Chem. 2011, 50, 9849–9863. 35. Stanciu, C.; Jones, M. E.; Fanwick, P. E.; Abu-Omar, M. M. J. Am. Chem. Soc. 2007, 129, 12400–12401. 36. Tsurugi, H.; Saito, T.; Tanahashi, H.; Arnold, J.; Mashima, K. J. Am. Chem. Soc. 2011, 133, 18673–18683. 37. Nishiyama, H.; Ikeda, H.; Saito, T.; Kriegel, B.; Tsurugi, H.; Arnold, J.; Mashima, K. J. Am. Chem. Soc. 2017, 139, 6494–6505. 38. Pintauer, T.; Matyjaszewski, K. Chem. Soc. Rev. 2008, 37, 1087–1097. 39. Klein Gebbink, R. J. M., Moret, M., Eds.; In Non-Noble Metal Catalysis: Molecular Approaches and Reactions; Wiley-VCH, 2019. 40. Bullock, R. M.; Chen, J. G.; Gagliardi, L.; Chirik, P. J.; Farha, O. K.; Hendon, C. H.; Jones, C. W.; Keith, J. A.; Klosin, J.; Minteer, S. D.; Morris, R. H.; Radosevich, A. T.; Rauchfuss, T. B.; Strotman, N. A.; Vojvodic, A.; Ward, T. R.; Yang, J. Y.; Surendranath, Y. Science 2020, 369, eabc3183. 41. Gibson, V. C.; Redshaw, C.; Solan, G. A. Chem. Rev. 2007, 107, 1745–1776. 42. Britovsek, G. J. P.; Gibson, V. C.; Kimberley, B. S.; Maddox, P. J.; McTavish, S. J.; Solan, G. A.; White, A. J. P.; Williams, D. J. Chem. Commun. 1998, 7, 849–850. 43. (a) Small, B. L.; Brookhart, M.; Bennett, A. M. A. J. Am. Chem. Soc. 1998, 120, 4049–4050; (b) Small, B. L.; Brookhart, M. J. Am. Chem. Soc. 1998, 120, 7143–7144. 44. (a) De Bruin, B.; Bill, E.; Bothe, E.; Weyhermüller, T.; Wieghardt, K. Inorg. Chem. 2000, 39, 2936–2947; (b) Budzelaar, P. H. M.; De Bruin, B.; Gal, A. W.; Wieghardt, K.; Van Lenthe, J. H. Inorg. Chem. 2001, 40, 4649–4655; (c) Enright, D.; Gambarotta, S.; Yap, G. P. A.; Budzelaar, P. H. M. Angew. Chemie - Int. Ed. 2002, 41, 3873–3876. 45. Bart, S. C.; Lobkovsky, E.; Chirik, P. J. J. Am. Chem. Soc. 2004, 126, 13794–13807. 46. (a) Bart, S. C.; Chłopek, K.; Bill, E.; Bouwkamp, M. W.; Lobkovsky, E.; Neese, F.; Wieghardt, K.; Chirik, P. J. J. Am. Chem. Soc. 2006, 128, 13901–13912; (b) Wile, B. M.; Trovitch, R. J.; Bart, S. C.; Tondreau, A. M.; Lobkovsky, E.; Milsmann, C.; Bill, E.; Wieghardt, K.; Chirik, P. J. Inorg. Chem. 2009, 48, 4190–4200; (c) Stieber, S. C. E.; Milsmann, C.; Hoyt, J. M.; Turner, Z. R.; Finkelstein, K. D.; Wieghardt, K.; Debeer, S.; Chirik, P. J. Inorg. Chem. 2012, 51, 3770–3785. 47. Chirik, P. J. Acc. Chem. Res. 2015, 48, 1687–1695. 48. Bouwkamp, M. W.; Bowman, A. C.; Lobkovsky, E.; Chirik, P. J. J. Am. Chem. Soc. 2006, 128, 13340–13341.

Redox-Active Ligands in Organometallic Chemistry

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. 97. 98. 99. 100. 101. 102. 103. 104. 105. 106.

441

Hu, L.; Chen, H. J. Am. Chem. Soc. 2017, 139, 15564–15567. Joannou, M. V.; Hoyt, J. M.; Chirik, P. J. J. Am. Chem. Soc. 2020, 142, 5314–5330. Zhou, Y.-Y.; Hartline, D. R.; Steiman, T. J.; Fanwick, P. E.; Uyeda, C. Inorg. Chem. 2014, 53, 11770–11777. (a) Rounds, H. R.; Zeller, M.; Uyeda, C. Organometallics 2018, 37, 545–550; (b) Hartline, D. R.; Zeller, M.; Uyeda, C. J. Am. Chem. Soc. 2017, 139, 13672–13675. Maity, A. K.; Zeller, M.; Uyeda, C. Organometallics 2018, 37, 2437–2441. Adolph, C. M.; Lee, S. A.; Zeller, M.; Uyeda, C. Tetrahedron 2019, 75, 3336–3340. Powers, I. G.; Andjaba, J. M.; Luo, X.; Mei, J.; Uyeda, C. J. Am. Chem. Soc. 2018, 140, 4110–4118. Zhou, Y. Y.; Uyeda, C. Science 2019, 363, 857–862. Mosmann, T. J. Immunol. 1983, 65, 55–63. Chang, M. C.; Dann, T.; Day, D. P.; Lutz, M.; Wildgoose, G. G.; Otten, E. Angew. Chemie - Int. Ed. 2014, 53, 4118–4122. Travieso-Puente, R.; Broekman, J. O. P.; Chang, M.-C.; Demeshko, S.; Meyer, F.; Otten, E. J. Am. Chem. Soc. 2016, 138, 5503–5506. Broere, D. L. J.; Mercado, B. Q.; Lukens, J. T.; Vilbert, A. C.; Banerjee, G.; Lant, H. M. C.; Lee, S. H.; Bill, E.; Sproules, S.; Lancaster, K. M.; Holland, P. L. Chem. Eur. J. 2018, 24, 9417–9425. Broere, D. L. J.; Mercado, B. Q.; Bill, E.; Lancaster, K. M.; Sproules, S.; Holland, P. L. Inorg. Chem. 2018, 57, 9580–9591. Chang, M.-C.; Roewen, P.; Travieso-Puente, R.; Lutz, M.; Otten, E. Inorg. Chem. 2015, 54, 379–388. Chang, M.-C.; Otten, E. Inorg. Chem. 2015, 54, 8656–8664. Smith, A. L.; Hardcastle, K. I.; Soper, J. D. J. Am. Chem. Soc. 2010, 132, 14358–14360. Lippert, C. A.; Amstein, S. A.; David Sherrill, C.; Soper, J. D. J. Am. Chem. Soc. 2010, 132, 3879–3892. Mukherjee, C.; Weyhermüller, T.; Bothe, E.; Chaudhuri, P. Inorg. Chem. 2008, 47, 2740–2746. Jacquet, J.; Salanouve, E.; Orio, M.; Vezin, H.; Blanchard, S.; Derat, E.; Desage-El Murr, M.; Fensterbank, L. Chem. Commun. 2014, 50, 10394–10397. Jacquet, J.; Chaumont, P.; Gontard, G.; Orio, M.; Vezin, H.; Blanchard, S.; Desage-El Murr, M.; Fensterbank, L. Angew. Chemie - Int. Ed. 2016, 55, 10712–10716. Jacquet, J.; Blanchard, S.; Derat, E.; Desage-El Murr, M.; Fensterbank, L. Chem. Sci. 2016, 7, 2030–2036. Jacquet, J.; Cheaib, K.; Ren, Y.; Vezin, H.; Orio, M.; Blanchard, S.; Fensterbank, L.; Desage-El Murr, M. Chem. Eur. J. 2017, 23, 15030–15034. Kraft, S. J.; Williams, U. J.; Daly, S. R.; Schelter, E. J.; Kozimor, S. A.; Boland, K. S.; Kikkawa, J. M.; Forrest, W. P.; Christensen, C. N.; Schwarz, D. E.; Fanwick, P. E.; Clark, D. L.; Conradson, S. D.; Bart, S. C. Inorg. Chem. 2011, 50, 9838–9848. Schelter, E. J.; Wu, R.; Scott, B. L.; Thompson, J. D.; Cantat, T.; John, K. D.; Batista, E. R.; Morris, E. D.; Kiplinger, J. L. Inorg. Chem. 2010, 49, 924–933. Kraft, S. J.; Fanwick, P. E.; Bart, S. C. Inorg. Chem. 2010, 49, 1103–1110. Zi, G.; Blosch, L. L.; Jia, L.; Andersen, R. A. Organometallics 2005, 24, 4602–4612. Zi, G.; Jia, L.; Werkema, E. L.; Walter, J. P.; Gottfriedsen, J. P.; Andersen, R. A. Organometallics 2005, 24, 4251–4264. Diaconescu, P. L.; Arnold, P. L.; Baker, T. A.; Mindiola, D. J.; Cummins, C. C. J. Am. Chem. Soc. 2000, 122, 6108–6109. Evans, W. J.; Traina, C. A.; Ziller, J. W. J. Am. Chem. Soc. 2009, 131, 17473–17481. Matson, E. M.; Opperwall, S. R.; Fanwick, P. E.; Bart, S. C. Inorg. Chem. 2013, 52, 7295–7304. Matson, E. M.; Franke, S. M.; Anderson, N. H.; Cook, T. D.; Fanwick, P. E.; Bart, S. C. Organometallics 2014, 33, 1964–1971. Anderson, N. H.; Odoh, S. O.; Yao, Y.; Williams, U. J.; Schaefer, B. A.; Kiernicki, J. J.; Lewis, A. J.; Goshert, M. D.; Fanwick, P. E.; Schelter, E. J.; Walensky, J. R.; Gagliardi, L.; Bart, S. C. Nat. Chem. 2014, 6, 919–926. Kiernicki, J. J.; Fanwicka, P. E.; Bart, S. C. Chem. Commun. 2014, 50, 8189–8192. Kiernicki, J. J.; Cladis, D. P.; Fanwick, P. E.; Zeller, M.; Bart, S. C. J. Am. Chem. Soc. 2015, 137, 11115–11125. Coughlin, E. J.; Qiao, Y.; Lapsheva, E.; Zeller, M.; Schelter, E. J.; Bart, S. C. J. Am. Chem. Soc. 2019, 141, 1016–1026. Boyer, J. L.; Cundari, T. R.; Deyonker, N. J.; Rauchfuss, T. B.; Wilson, S. R. Inorg. Chem. 2009, 48, 638–645. Ringenberg, M. R.; Kokatam, S. L.; Heiden, Z. M.; Rauchfuss, T. B. J. Am. Chem. Soc. 2008, 130, 788–789. Ringenberg, M. R.; Nilges, M. J.; Rauchfuss, T. B.; Wilson, S. R. Organometallics 2010, 29, 1956–1965. Deibel, N.; Schweinfurth, D.; Hohloch, S.; Fiedler, J.; Sarkar, B. Chem. Commun. 2012, 48, 2388–2390. Bubrin, M.; Schweinfurth, D.; Ehret, F.; Záliš, S.; Kvapilová, H.; Fiedler, J.; Zeng, Q.; Hartl, F.; Kaim, W. Organometallics 2014, 33, 4973–4985. For examples see (a) Hennessy, E. T.; Betley, T. A. Science. 2013, 340, 591–595.(b) Iovan, D. A.; Wilding, M. J. T.; Baek, Y.; Hennessy, E. T.; Betley, T. A. Angew. Chem., Int. Ed. 2017, 56, 15599–15602; (c) Kuijpers, P. F.; Tiekink, M. J.; Breukelaar, W. B.; Broere, D. L. J.; van Leest, N. P.; van der Vlugt, J. I.; Reek, J. N. H.; de Bruin, B. Chem. Eur. J. 2017, 23, 7945–7952; (d) Zhou, M.; Lankelma, M.; van der Vlugt, J. I.; de Bruin, B. Angew. Chem. Int., Ed. 2020, 59, 11073–11079 Broere, D. L. J.; De Bruin, B.; Reek, J. N. H.; Lutz, M.; Dechert, S.; Van Der Vlugt, J. I. J. Am. Chem. Soc. 2014, 136, 11574–11577. Broere, D. L. J.; van Leest, N. P.; De Bruin, B.; Siegler, M. A.; Van Der Vlugt, J. I. Inorg. Chem. 2016, 55, 8603–8611. Bagh, B.; Broere, D. L. J.; Sinha, V.; Kuijpers, P. F.; Van Leest, N. P.; De Bruin, B.; Demeshko, S.; Siegler, M. A.; Van Der Vlugt, J. I. J. Am. Chem. Soc. 2017, 139, 5117–5124. Broere, D. L. J.; Metz, L. L.; De Bruin, B.; Reek, J. N. H.; Siegler, M. A.; Van Der Vlugt, J. I. Angew. Chemie - Int. Ed. 2015, 54, 1516–1520. Ren, Y.; Cheaib, K.; Jacquet, J.; Vezin, H.; Fensterbank, L.; Orio, M.; Blanchard, S.; Desage-El Murr, M. Chem. Eur. J. 2018, 24, 5086–5090. van Leest, N. P.; Tepaske, M. A.; Oudsen, J. P. H.; Venderbosch, B.; Rietdijk, N. R.; Siegler, M. A.; Tromp, M.; Van Der Vlugt, J. I.; De Bruin, B. J. Am. Chem. Soc. 2020, 142, 552–563. van Leest, N. P.; Tepaske, M. A.; Venderbosch, B.; Oudsen, J. H.; Tromp, M.; Van Der Vlugt, J. I.; De Bruin, B. ACS Catal. 2020, 10, 7449–7463. Whittaker, J. W. Chem. Rev. 2003, 103, 2347–2363. Chaudhuri, P.; Hess, M.; Müller, J.; Hildenbrand, K.; Bill, E.; Weyhermüller, T.; Wieghardt, K. J. Am. Chem. Soc. 1999, 121, 9599–9610. Wang, Y.; Stack, T. D. P. J. Am. Chem. Soc. 1996, 118, 13097–13098. Wang, Y.; DuBois, J. L.; Hedman, B.; Hodgson, K. O.; Stack, T. D. P. Science 1998, 279, 537–540. Königsmann, M.; Donati, N.; Stein, D.; Schönberg, H.; Harmer, J.; Sreekanth, A.; Grützmacher, H. Angew. Chemie - Int. Ed. 2007, 46, 3567–3570. Wang, K.; Stiefel, E. I. Science 2001, 291, 106–109. Harrison, D. J.; Nguyen, N.; Lough, A. J.; Fekl, U. J. Am. Chem. Soc. 2006, 128, 11026–11027. (a) Dang, L.; Shibl, M. F.; Yang, X.; Alak, A.; Harrison, D. J.; Fekl, U.; Brothers, E. N.; Hall, M. B. J. Am. Chem. Soc. 2012, 134, 4481–4484; (b) Dang, L.; Shibl, M. F.; Yang, X.; Harrison, D. J.; Alak, A.; Lough, A. J.; Fekl, U.; Brothers, E. N.; Hall, M. B. Inorg. Chem. 2013, 52, 3711–3723. Hirsekorn, K. F.; Hulley, E. B.; Wolczanski, P. T.; Cundari, T. R. J. Am. Chem. Soc. 2008, 130, 1183–1196. Shaik, S.; Chen, H.; Janardanan, D. Nat. Chem. 2011, 3, 19–27.

1.14 Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry Elizabeth T Papish, Sanjit Das, Weerachai Silprakob, Chance M Boudreaux, and Sonya Manafe, Department of Chemistry, The University of Alabama, Tuscaloosa, AL, United States © 2022 Elsevier Ltd. All rights reserved.

1.14.1 1.14.1.1 1.14.1.2

Introduction Inspiration from nature Which types of reactions are enhanced by metal-ligand bifunctional catalysts involving protic and hydrogen bonding ligands? 1.14.1.3 Hydrogenation considerations 1.14.1.4 Which metals are used frequently for hydrogenation catalysis? 1.14.1.5 Considering the pKa value of the pendant proton 1.14.1.6 Ligand rigidity, vulnerability to rearrangement, and ring size 1.14.1.7 List of criteria for evaluating catalysts for enhancements with protic groups or hydrogen bonding groups 1.14.1.8 Which metals are used frequently for MLBC? 1.14.1.9 Overview: Ligand architectures 1.14.2 Noyori’s catalyst and related systems with NdH near the metal center 1.14.2.1 Noyori’s catalyst and closely related Ru catalysts 1.14.2.2 Iridium catalysts with NH near the metal center 1.14.3 Shvo’s catalyst and related systems with OH groups on a cyclopentadienyl ring 1.14.3.1 Shvo’s catalyst and other closely related Ru catalysts 1.14.3.2 Iron analogues of Shvo’s catalyst 1.14.3.3 Other metals for Shvo catalyst analogues 1.14.4 Pyridinol derived metal complexes as catalysts for hydrogenation and other reactions 1.14.4.1 Motivation and background 1.14.4.2 Ligands containing one pyOH group 1.14.4.3 Chelating ligands containing two pyOH groups 1.14.4.3.1 Early work: CO2 hydrogenation and other reactions 1.14.4.3.2 Recent work: Hydrogenation, dehydrogenation and related reactions 1.14.4.3.3 CO2 hydrogenation and related reactions 1.14.4.3.4 Electrocatalytic oxidation and reduction reactions 1.14.4.3.5 Summary of pyridinol based ligands for organometallic catalysis 1.14.5 Tridentate facial arrangements of hydrogen bond donors or acceptors 1.14.6 Proton responsive pincer ligands 1.14.7 Perspective Acknowledgments References

443 444 444 445 446 446 447 447 447 448 448 448 449 450 450 451 452 453 453 454 455 455 456 457 459 460 460 464 467 467 468

Nomenclature 4,4ʹ-dhbp 6,6ʹ-dhbp bpy CH3CN CO CO2 Cp Cp CpOH CV Cy DFT Dhbp DMF ee EPR

442

4,4ʹ-Dihydroxy-2,2ʹ-bipyridine 6,6ʹ-Dihydroxy-2,2ʹ-bipyridine 2,2ʹ-Bipyridine Acetonitrile Carbon monoxide (as a free molecule) or carbonyl (as a ligand) Carbon dioxide Cyclopentadienyl anion ¼ [C5H5]− Pentamethylcyclopentadienyl anion ¼ [C5(CH3)5]− Hydroxycyclopentadienyl anion ¼ [C5(OH)H4]− Cyclic voltammetry Cyclohexyl Density functional theory Dihydroxy-2,2ʹ-bipyridine N,N-dimethylformamide Enantiomeric excess Electron paramagnetic resonance

Comprehensive Organometallic Chemistry IV

https://doi.org/10.1016/B978-0-12-820206-7.00080-9

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

Et [HCO2]− HDO KIE LRPHT Me MLBC MLC MST NHC OEC OH OMe pyOH THF Tren Ttz TOF TON

1.14.1

443

Ethyl Formate Hydrodeoxygenation Kinetic isotope effect Ligand rotation-promoted hydrogen transfer pathway Methyl Metal ligand bifunctional catalysis Metal ligand cooperation N,N0 ,N00 -[2,20 ,200 -nitrilotris(ethane-2,1-diyl)]tris(2,4,6-trimethylbenzenesulfonamido) N-heterocyclic carbene Oxygen evolving complex Hydroxy Methoxy 2-Pyridinol Tetrahydrofuran Tris(2-aminoethyl)amine Tris(1,2,4-triazolyl)borate anion Turnover frequency Turnover number

Introduction

Proton responsive ligands are frequently used in organometallic catalysis reactions in order to control the delivery of protons near the metal site. This type of catalysis is termed “metal ligand bifunctional catalysis” (MLBC) or “metal ligand cooperation” (MLC).1–5 Traditional catalysis usually features minimal interaction between the ligand(s) and the substrate, and the main purpose of the ligands is to provide steric and electronic interactions with the metal (Fig. 1A). In contrast, metal ligand bifunctional catalysts offer hydrogen bond donors or acceptors which can interact directly with the substrate (Fig. 1B).6,7 Typically these interactions serve to provide a pathway with lowered activation energy to accelerate catalytic transformations which convert the substrate to product. Additionally, the catalyst can sometimes be “switched” between active and inactive forms by protonation events; these are proton responsive catalysts.8 These interactions are not limited to hydrogen bonds and can also include other electrostatic interactions that impact catalysis, most commonly from fluorinated organic ligands or pendant ammonium groups,9 but hydrogen bonding and protic ligands are the main focus herein. Inherent in this type of catalysis is a requirement that the metal chosen is not sensitive to protons. Furthermore, the term “organometallic” refers to the presence of metal-carbon bonds, and while alkyl groups also provide carbon donors, metal alkyl complexes are typically sensitive to protons and therefore are not used with proton bearing ligands and are not discussed herein. Instead, the carbon donors in proton-responsive systems are more commonly cyclopentadienyl (Cp), CO, and carbene donor ligands.10 These are all strong field ligands and other isoelectronic ligands are frequently included in the field of organometallic chemistry. For this reason, several N (e.g. 2,2ʹ-bipyridine (bpy)) and P (e.g. phosphine) donor ligands are included here. Organometallic catalysis also includes transformations at carbon atoms that are catalyzed by metals, including hydrogenation and C-H activation reactions.

Fig. 1 (A) Traditional metal-based catalysis in which the ligand does not interact with the substrate. (B) Metal ligand bifunctional catalysis most often includes hydrogen bond donors and/or hydrogen bond acceptors on the ligand which interact with the substrate to provide a pathway for catalysis with a lower activation energy. X is typically an electronegative element (X ¼ O, N, etc.).

444

1.14.1.1

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

Inspiration from nature

Several metalloenzyme active sites feature hydrogen bonding groups near the metal center which serve to enhance rates of proton transfer events (Fig. 2). For example, the mono-iron hydrogenase enzyme (2a) uses H2 as a source of hydride for transfer to a cationic substrate, methenyltetrahydromethanopterin.4,11 H2 is cleaved heterolytically by placing H+ on the ligand O− group and hydride is placed on the Fe center (L ¼ H−) in 2a. Thus, 2a is generated which can then transfer hydride to substrate. Herein, a basic group on the ligand in the form of a pyridinolate ring is essential for the facile splitting of H2 as H+ and H−. Likewise, the dangling NH group above the active site in the di-iron hydrogenase (2b) allows dihydrogen to be split as H− (on Fe) and H+ on the ligand, forming an R2NH+2 group.12 Both experimental and computational studies using both enzymes and model compounds suggest that the pendant basic group in 2b is critical for its activity.13,14 Many organometallic complexes and catalysts that mimic both the structure and function of 2b have been reported and these are described further herein.15–19 Both of these enzymes perform proton reduction and the reverse reaction, H2 oxidation (Fig. 2B). Several other metalloenzymes (e.g. carbon monoxide dehydrogenase, lactate racemase, alcohol dehydrogenase, and others) also feature hydrogen bonding interactions that are critical to rate acceleration.4 For more in depth coverage of this topic, the reader is directed to several review articles.4,20 Herein, the focus is on using these enzyme active site structures as inspiration for small molecule catalysts, which act as structural and functional mimics to reproduce both the hydrogen bonding interactions and the catalytic rate accelerations.

1.14.1.2 Which types of reactions are enhanced by metal-ligand bifunctional catalysts involving protic and hydrogen bonding ligands? Frequently, reactions involving dihydrogen are accelerated by metal catalysts that can accept or donate protons. This includes both proton reduction and the reverse reaction, H2 oxidation, as illustrated in Fig. 2.6,20 Mimicking the hydrogenase enzymes, several organometallic catalysts have been devised using both pyridinol ligands (like 2a)3,8 and pendant amines (like 2b),6 and these are covered in greater detail in this chapter. Typically, the role of ligands which can accept or donate protons is to provide a lower energy pathway for proton transfer events.6,21 Similarly, these sorts of ligands are also frequently used for CO2 reduction (Fig. 3), which can

Fig. 2 (A) Mono-iron hydrogenase (2a) and di-iron hydrogenase. (2b) active sites. (B) Schemes for the reactions catalyzed by 2a and 2b. 2a catalyzes the transfer of hydride to methenyltetrahydromethanopterin and 2b catalyzes dihydrogen oxidation and the reverse reaction, proton reduction.

Fig. 3 CO2 hydrogenation and reduction reactions are often accelerated proton responsive metal catalysts.

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

445

achieved by either electrochemical or photochemical methods.22 Carbon dioxide reduction can form CO, formic acid, formate, carbonate, or other products involving multiple electrons and protons (e.g. formaldehyde, methanol, methane, etc.).23–27 Related to this process is CO2 hydrogenation, in which H2 is used as the reducing agent rather than protons and electrons (Fig. 3).23 Furthermore, other closely related reactions are frequently covered in this chapter, including hydrogenation of ketones and aldehydes to form alcohols and the reverse reaction which is acceptorless alcohol oxidation which releases H2.3

1.14.1.3

Hydrogenation considerations

Proton responsive ligands bound to ruthenium and iridium have led to exceptionally active catalysts for various processes, with hydrogenation reactions being especially prominent. Principally, these studies have used Ru(II) and Ir(III) which are d6 metals and are especially suitable for hydrogen activation. Typically, a five coordinate metal binds H2 via an empty orbital (Fig. 4A), forming a six coordinate dihydrogen metal complex, as shown in Fig. 4B. When H2 binds to the metal in octahedral or pseudo octahedral symmetry, an empty eg orbital accepts electron density from H2 while the t2g set of orbitals are occupied and one t2g orbital can donate electron density towards the s orbital of H2 (Fig. 4A).28,29 This interaction serves to cleave the HdH bond and leads to either a dihydride complex (homolytic HdH bond cleavage) or a hydride and a proton located on the ligand or the base/solvent (heterolytic HdH bond cleavage) (Fig. 4B). Homolytic HdH bond cleavage (Fig. 4B) is accompanied by an increase in formal oxidation state (Mn goes to Mn+2) at the metal center as electron density is transferred from the metal to the H2 to form the hydride ligands. At the same time, the metal becomes a seven coordinate dihydride complex (or more generally, the coordination number increases by two). Thus, for homolytic HdH bond cleavage to occur, the metal must be able to accommodate an increase in coordination number (via gaining two bonds or the metal can reduce its coordination number by subsequent ligand loss) as well as a higher oxidation state. For these reasons, homolytic HdH bond cleavage tends to be more rare vs heterolytic HdH bond cleavage, and homolytic bond cleavage typically occurs for larger transition metals that can accommodate the increase in coordination number.29,30 From an electron counting perspective, both the dihydrogen complex and the dihydride have the same electron count; e.g. with Ru(II) as the metal the 18e− dihydrogen complex becomes an 18e− Ru(IV) dihydride in Fig. 4B. Heterolytic HdH bond cleavage (Fig. 4B) can involve proton transfer from an acidic dihydrogen complex28,30 to a solvent or base molecule. This generates a six-coordinate complex which has a −1 charge relative to a neutral and five coordinate (ML5) starting material. Commonly within this book chapter, the ligands contain basic sites and heterolytic HdH bond cleavage occurs with H+ on the ligand and H− on the metal. Thus, neither the charge on the metal complex nor the oxidation number of the metal changes. However, in the course of this reaction, the ligand is protonated and may become a less electron donating ligand if the protonation site is close enough to the metal and conjugated to influence the electron donor properties of the ligand. The factors that influence which pathway (Fig. 4B) is followed include: (1) Is the dihydrogen complex sufficiently acidic? (2) Is the base, solvent, or the internal basic site on the ligand sufficiently basic to accept the proton?, and (3) Are there considerations that favor the homolytic pathway such as a metal that prefers a high coordination number and a high oxidation state (e.g. third row transition metals such as Os, W, etc. often favor homolytic cleavage).29–33

Fig. 4 (A) Interactions between H2 molecular orbitals and d6 octahedral metal complexes (e.g. Ru(ll), Ir(lll), etc.). These interactions explain why low spin d6 metal complexes are especially good at activating H2 and small molecules containing strong sigma bonds. (B) Homolytic vs heterolytic H2 activation.

446

1.14.1.4

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

Which metals are used frequently for hydrogenation catalysis?

As mentioned above, d6 low spin metals are well suited for hydrogenation catalysis due to orbital considerations (Fig. 4A). Additionally, some hydrogenation reactions proceed via oxidative addition (Fig. 4B) and thus favor metals with a tendency to perform 2e− processes rather than 1e− processes involving radicals. Given these trends, Ru(II), Ir(III) and other second and third row transition metal complexes are often used as hydrogenation catalysts, but with the right supporting ligand set, it is possible for first row metals (e.g. iron, cobalt)34,35 to be used.36,37 Typically the ligands are strong field to favor a low spin configuration1,38 and in some cases the ligands can accommodate unpaired electron density (2e− processes can be shared as 1e− transfer at the metal and 1e− at the ligand).36 Furthermore, strong field ligands and/or second and third row transition metals lead to slow ligand exchange rates such that the supporting ligands are not exchanged readily in the course of catalysis.39 The first ligands used to make highly active metal-ligand bifunctional hydrogenation catalysts with Ru included the Noyori system2,40,41 and the Shvo system (Fig. 5, 5a and 5b, respectively).42–44 The Noyori catalyst (5a) bears resemblance to the di-iron hydrogenase (Fig. 2, 2b) in that both catalysts bear NH groups near the metal center. More recently, pyridinol based ligands have been frequently used with Ru and Ir (Fig. 5, e.g. 5c) to construct highly active hydrogenation catalysts.3,8,45 Similarly, 5b and 5c feature protic OH groups near the metal center to mimic 2a (Fig. 2), the mono-iron hydrogenase enzyme.20

1.14.1.5

Considering the pKa value of the pendant proton

Typically, the pKa value of the pendant protic group has been around 4–10 in water in order to serve as both a proton donor and a base. For example, the pKa of the OH groups in 5c is 4.1 in water, with simultaneous removal of both protons around pH 4.46 This allows for the OH groups to be deprotonated at neutral pH and for the O− groups on the ligand to accept protons as a pendant (weak) base. We note that for ligands like 6,6ʹ-dihydroxybipyridine (6,6ʹ-dhbp) in 5c, the pKa of the free organic ligand is estimated at 8.5 in water.47,48 Metal coordination provides a nearby Lewis acid to remove electron density from the organic ligand, and thus the resulting pKa of the metal complex is typically 3–4 pH units lower.47,49,50 Of course, if the pendant OH or NH group is further from the metal center and/or not conjugated to that metal center, then the impact of metal coordination on the resulting pKa values is much less. Although the pKa of the OH group within the monoiron hydrogenase 2a has not been measured directly, it is likely to be similar to Ru(II) complexes of 6,6ʹ-dhbp (pKa 5).51 Similarly, for the diiron hydrogenase (2b) the pKa value of the pendant NH2 group (formed by protonation of 2b) is estimated as 7.2 in the reduced state and 100 articles), orange, moderately used (in 10–100 articles), yellow, used rarely (in 1–10 articles), gray, not used.

With group 7 metals, Mn and Re complexes have incorporated protic groups and the main application has been CO2 reduction chemistry using planar bidentate and tridentate ligands50,71 and oxygen activation using tripodal ligands.5 The results of these studies are described herein. The radioactive nature of Tc has prevented extensive organometallic studies. Group 6 metal complexes bearing proton responsive ligands have been limited but selected examples are included herein.72,73 The early transition metals (groups 3–5) are very oxophilic and when they form organometallic complexes with metal-carbon bonds they are very sensitive to trace moisture and other proton sources. Likewise, the group 12 organometallic complexes74–76 tend to be moisture sensitive and incompatible with proton sources, including from a proton responsive ligand. While there are several zinc coordination complexes that feature hydrogen bonds near the metal center,74,77–79 these are beyond the scope of this chapter and will only be mentioned briefly.

1.14.1.9

Overview: Ligand architectures

This chapter will be organized by ligand type with various metal complexes and various types of catalysis being discussed within each ligand type. Noyori’s catalyst and other ligands that offer NH groups near the metal center are discussed first (Section 1.14.2).2,80 In Section 1.14.3, the Shvo catalyst positions OH groups on the ligand scaffold via a CpOH ligand.42,54,65 In Section 1.14.4, pyridinol based metal complexes are described. Like the Shvo system, the pyridinol ligand places OH groups near the metal center for transfer of H+ to substrates.8,20 Other ligands that offer OH groups near the metal center (e.g. porphyrins and bpy ligands with dangling OH groups)22,71 are also discussed here due to a similarity in structure and a similarity in reactivity. Section 1.14.5 switches gears with an emphasis on tripodal ligands which donate NH, OH or other hydrogen bond acceptors.5 In this section, we also mention other ligands containing hydrogen bond acceptors, including bidentate NN and PP ligands studied by Periana and DuBois,6,68 as a point of comparison. Section 1.14.6 focuses on pincer ligands1 with proximal or remote protic groups.

1.14.2

Noyori’s catalyst and related systems with NdH near the metal center

1.14.2.1

Noyori’s catalyst and closely related Ru catalysts

Noyori’s catalyst (5a, Figs. 5 and 7) and the related Noyori-Ikariya catalyst (Fig. 7, 7a) are highly active and enantioselective catalysts that have streamlined the industrial synthesis of chiral alcohols; this work was honored with the Nobel Prize in 2001.40,41,80–86 Another highly active catalyst from Ohkuma (Fig. 7, 7b) has a similar structure to 5a but contains a chelating aryl ring (from a facial tridentate NNC donor).87 This catalyst achieves >99% ee for ketone hydrogenation with a turnover frequency of 35,000 min−1 at a mild pressure of 50 atm H2. Several other derivatives of 5a and 7a have been reported which are reviewed elsewhere.2,64,88 Similar catalysts using Ir will be discussed below. These catalysts were proposed to transfer hydrogen to substrates via a concerted pathway involving a six-membered transition state with the NH groups transferring protons to the substrate (Fig. 8, 8a). Computations in the gas phase were used to support this mechanism along with the fact that removal of the NH groups harms catalyst performance. However, more recently, experimental and computational data supports a different mechanism.89–91 There were several experimental inconsistencies with the concerted mechanism, reported as early as 2001.56,89–91 Dub and Gordon have demonstrated by computational studies that this mechanism is not viable in the solution phase and the hydrogenation of substrates occurs by a step-wise pathway with NdH or NdK (K arises from the base, KOtBu) stabilizing the bound substrate near the metal center (Fig. 8, 8b). Thus while NH (or NK) interactions with the substrate are key to lowering the activation barrier, the source of H+ is most likely a metal bound dihydrogen complex (8d) rather than the NH groups.2,57

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

449

Fig. 7 Highly active and enantioselective MLC catalysts using Ru and Ir (5a, 7a–7e). Also shown are highly chemoselective catalysts 7f and 7g.

Fig. 8 (A) A concerted mechanism for hydrogenation by Noyori’s catalyst is no longer accepted. While this transition state was located in the gas phase, it is not located in solution computations. (B) The computational evidence supports a stepwise MLC mechanism wherein hydride is transferred prior to a proton being transferred to substrate from a metal dihydrogen complex. X ¼ H or K (K present at high concentration of base, KOtBu).

1.14.2.2

Iridium catalysts with NH near the metal center

Similarly, iridium catalysts (Fig. 7) bearing NH groups accomplish the rapid and enantioselective hydrogenation of ketones. Catalyst 7c (Fig. 7) was synthesized first featuring a chiral bidentate PN ligand.92 However, this catalyst was vulnerable to side reactions including the formation of an inactive byproduct formulated as [(PN)2IrH2]+.93 Aiming to avoid this side product as well as the formation of inactive dimers and trimers, Xie and Zhou designed a catalyst (7d, Fig. 7) featuring a bulky tridentate ligand. Catalyst 7d hydrogenates acetophenone with substrate:catalyst ratios of up to 5 x 106 and 98% ee.94 Catalyst 7e (Fig. 7) is also highly active and achieves up to 99.9% ee.95 The success of catalysts 7d–7e suggest that there is an advantage to using tridentate ligands to supply the NH group because they block the formation of dimers, trimers, and other inactive species. Furthermore, Dub, Gordon, and coworkers have developed NNS donor ligands with NH groups near the metal center.96 These ligands have formed iridium and ruthenium complexes (7f and 7g, Fig. 7) that perform the selective hydrogenation of methyl trifluoroacetate under mild conditions with turnover numbers as high as 10,000 for the best iridium catalysts. Again, these studies show that tridentate and tetradentate ligands appear to be advantageous over bidentate ligands, especially with iridium, for the reasons discussed above.

450

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

1.14.3

Shvo’s catalyst and related systems with OH groups on a cyclopentadienyl ring

1.14.3.1

Shvo’s catalyst and other closely related Ru catalysts

The Shvo catalyst (Figs. 5 and 9, structures 9a and 5b) was reported in 1985 and shown to be a highly active precatalyst for (de) hydrogenation reactions. The scope of catalytic reactions includes transfer hydrogenation of imines and carbonyl compounds, direct hydrogenation of substrates, and amine and alcohol oxidation.1,43,97,98 Recent articles have expanded the scope of the Shvo catalyst to include reactions involving hydrogen removal including dehydrogenation of diols,99 N-formylation of lactams,100 and decarboxylative N-alkylation using a hydrogen-borrowing approach (Fig. 10C).101 The hydrogen-borrowing approach is illustrated in Fig. 10, and it generally involves starting with an alcohol dehydrogenation reaction to generate a more reactive ketone or aldehyde intermediate that can be functionalized and then re-hydrogenated.102 Frequently, a hydrogen-borrowing approach is used for the amination and for the b-alkylation of alcohols (Fig. 10A and B, respectively).103,104 Furthermore asymmetric (de)hydrogenation reactions have been developed105–107 along with reactions relevant to biomass conversion including the hydrogenation of levulinic acid derived from cellulose to form valerolactone.108,109 Shvo’s catalyst has been used in tandem reactions, in surface attached catalysts, and for polymer synthesis and other applications.110–117 Many new derivatives of the Shvo system have been developed.118–120 Catalysts similar to the Shvo system but using Fe and Ir as the metals will be discussed below. The Shvo catalyst uses a CpOH ligand (9e, Fig. 9) which is an anionic six electron p donor to most metal centers. Thus, CpOH is an L2X type ligand just like the Cp anion. However, the OH group serves as a reservoir for providing protons to substrates, and upon

Fig. 9 (A) The Shvo catalyst on the left (9a) dissociates in solution to produce 5b and 9c. This catalyst can hydrogenate ketones and aldehydes, and one proposed transition state is shown in 9d. (B) Organic CpOH ligands within the Shvo catalyst. Deprotonation typically liberates two electrons which are used to reduce the metal from Ru(II) to Ru(0).

Fig. 10 A hydrogen-borrowing approach to functionalizing alcohols. The alcohol is first dehydrogenated to activate it and produce an aldehyde or ketone. (A) This ketone or aldehyde can then react with an amine to form an imine which upon hydrogenation gives an amine. This results in amination of the starting alcohol. (B) Alternatively, the ketone can react with base and an electrophile to produce a functionalized ketone which can then be hydrogenated to produce a functionalized alcohol. Often, this last reaction is an alkylation reaction that proceeds via an aldol condensation. (C) Decarboxylative N-alkylation is shown.

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

451

donation of a proton the ligand becomes a neutral four electron donor (9f). Typically, two electrons are transferred to the metal upon proton donation, and for example Ru(II) becomes Ru(0) (e.g. 9c) after the CpOH ligand loses a proton. Thus, as shown in Fig. 9, the Ru complexes 5b (RuII, d6) and 9c (Ru0, d8) are 18 and 16 electron complexes, respectively. Complex 9a is considered the catalyst resting state and has been shown to dissociate rapidly in solution to generate 5b and “9c.”121 Structure 5b is believed to be the active catalyst which can transfer hydrogen to/from substrates cooperatively as H+ from the ligand and H− from the metal (via transition state 9d).42,122,123 Following the reported catalysis by Shvo et al., several mechanistic studies were performed by Casey, Guan, et al. to establish by kinetic isotope effects (KIE) studies that the reaction involves RudH and OdH bond cleavage to transfer H2 to substrates (e.g. aldehydes) by an outer sphere mechanism.124,125 This cooperative mechanism is also supported by studies of an dNHPh substituted Cp ring which is not sufficiently acidic to transfer H+ from the amine group and gave slow hydrogenation of benzaldehyde. However, catalysis occurs quickly if the NH group is protonated to form the more acidic dNH2Ph+ substituent.126 Studies of catalyst 5b have shown that it does not lose H2 directly because the distance between OH and RudH is too great and transition state is too high in energy; rather a solvent assisted loss of H2 is proposed (Fig. 11, via 11d).2,127 Recently, experimental and computational studies have shown that the 16 electron species 9c (Fig. 9) is not present to any significant extent in solution during (de)hydrogenation reactions involving organic alcohols and carbonyl compounds.128 Rather, 18 electron species involving a coordinated alkoxide group are favored and these species along with the hydride 5b participate in the catalysis (Fig. 11). Part of the evidence for this alternative mechanism included reports that pyridine, which is typically used as a catalyst poison in probing for nanoparticle formation, actually enhances the rate of catalysis which conflicts with the expected pyridine inhibition if production of 9c were a rate-limiting step.129 Pyridine perhaps serves the role of stabilizing a catalyst resting state in order to prevent catalyst decomposition.129 Overall, these studies suggest that the path forward in enhancing the properties of the Shvo catalyst is not to bulk up the system in order to stabilize the 16 electron species 9c, but rather to ensure via intermediate bulkiness on the CpOH ring that 5b can undergo transformations to 11a, 11d, and the other alkoxide bound species in Fig. 11 with little activation barrier for H2 loss or addition.128

1.14.3.2

Iron analogues of Shvo’s catalyst

Given the success of the Shvo catalyst, there has been substantial interest in incorporating an earth abundant metal in place of ruthenium. The use of iron has been most successful with the Knölker complex (12a) shown in Fig. 12.130 As reported by Casey and Guan in 2007, this Fe(II) complex is capable of hydrogenating ketones, aldehydes, and imines with H2 or isopropanol as the hydrogen source.131 Since then the scope has been expanded to include dehydrogenation of alcohols (with 13b, Fig. 13),132 hydrogenation of levulinic acid followed by cyclization to produce g - valerolactone (12e),108 decarboxylative N-alkylation (Fig. 10C) to construct heterocycles (12e),101 amination of alcohols by a hydrogen-borrowing approach (12e),104,133,134 tandem catalysis,135 and CdH activation.136 It is thought that greater steric bulk is needed for the Fe catalysts as compared to the Ru analogs to prevent the formation of decomposition products, hence catalysts like 12f were synthesized and studied.137 Mechanistic studies with Fe have been more limited than for Ru, but generally a sequence of steps involving Fe species analogous to 5b, 9c, and 9d (Fig. 9) for C]O hydrogenation has been proposed based on computational studies.138,139 To the best of our knowledge, a sequence of steps analogous to Fig. 11 has not been studied for C]O hydrogenation with iron. Complex 12a (which is a specific example of type 13a) was sensitive to air and light, and so others frequently used more stable analogs (13b and 13d) as entry points for catalysis (Fig. 13). Fe(0) nitrile complexes (12e and 13d) serve as precatalysts and were first developed by Funk and coworkers140 and later studied further by Zhao and coworkers.104,133 Catalyst 12e (an example of 13d) is a stable complex that can be thermally activated to generate an active catalyst for the (de)hydrogenation of aliphatic and aromatic ketones and imines. Here the active

Fig. 11 A revised mechanism for (de)hydrogenation reactions with the Shvo catalyst.

452

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

Fig. 12 Fe derivatives of the Shvo catalyst.

Fig. 13 Transformations involving the Knölker catalyst.

catalyst is proposed to be 13c, but it may also be a complex analogous to 13c with solvent or substrate weakly bound in the vacant site. Several other catalysts of this type have been developed and for most of these a role of the OH/OE substituent in facilitating (de)hydrogenation transformations is proposed (Fig. 12).54,137,141–144 This topic has been reviewed recently.44,65 Significantly, asymmetric Fe catalysts have been designed using the bifunctional OH-substituted cyclopentadienyl ligands with examples from Berkessel (e.g. 12h) and others.145 Efforts to induce chirality in the products by inclusion of chiral ligands replacing one CO have been more successful than the inclusion of chiral Cp-OH derived ligands.146 Catalyst 12h was typically generated in situ by CO substitution with the chiral ligand. Acetophenone was hydrogenated with moderate enantioselectivity (up to 31% enantiomeric excess). The asymmetric induction of chirality was improved greatly by using a chiral phosphoric acid ligand with 12g which gave up to 74% ee and 90% conversion in imine hydrogenation by Beller et al.147,148 Computational studies for asymmetric imine hydrogenation with 12g have shown that a stepwise mechanism operates for proton and hydride transfer and the catalyst resting state involves coordination of the deprotonated phosphoric acid to iron.149 This parallels the mechanism shown for Ru analogs in Fig. 11 in some ways. Others have used Fe catalysts for hydrogenation and achieved enantiomeric excess by using dynamic kinetic resolution with enzymes.150,151

1.14.3.3

Other metals for Shvo catalyst analogues

Considering group 9 analogs, cobalt complexes have been synthesized incorporating Cp-OH ligands, but thus far catalytic applications have not been pursued with these complexes.152,153 Rhodium and iridium complexes (14a) have been synthesized and used for the catalytic acceptorless dehydrogenation and formation of C]C bonds with unusual functional group tolerance (Fig. 14).154 The reaction was tolerant of ester and amide functional groups proximate to the formed double bond. Here again, the proximate OH group was proposed to play a role in H2 loss. Catalyst 14a1 was also used for the hydrogenolysis of arenols and aryl methyl ethers.155 This reaction is similar to hydrodeoxygenation reactions that will be discussed further below. Retro hydroformylation reactions were also performed with similar catalysts (14a–14c with the best results for 14c) in which aldehydes underwent CO loss followed by dehydrogenation to produce alkenes.156 Dehydrogenation of other CdC single bonds to produce conjugated double bonds with Ir catalyst 14a–c was also pursued, but the reaction gave higher percent conversion with [Cp IrCl2]2 vs the OH

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

453

Fig. 14 Iridium and rhodium complexes of OH-substituted Cp ligands.

Fig. 15 Platinum and rhenium complexes bearing OH-substituted Cp ligands.

substituted catalyst.157 A recent example of reactivity reversal was reported in which BdH bonds (e.g. from HBpin) are activated by 14d to produce 14e, with a B anion and an OH group.158 Typically BdH bond cleavage would give an OdB bond and a metal hydride. Catalytic applications have not yet been pursued with 14d or 14e, but they are a promising area for future research. The use of other metals with CpdOH ligands has been more limited. Recently a Pt complex (15a) that undergoes reductive elimination from Pt(IV) to Pt(II) with CdH bond formation was reported.159 The methyl group bound to Pt forms a CdH bond with the proton of the OH bearing ligand to form methane stoichiometrically (Fig. 15). Rhenium complexes bearing CpdOH ligands have also been reported (15b) and have been used as catalysts for the transfer hydrogenation of ketones and imines.160,161

1.14.4

Pyridinol derived metal complexes as catalysts for hydrogenation and other reactions

1.14.4.1

Motivation and background

As mentioned in the introduction, proton responsive ligands containing the pyridinol motif occur naturally in the mono-iron hydrogenase active site (2a, Figs. 2 and 16). The mono-iron hydrogenase enzyme features an OH group near the metal center19,20,162–165 that can be used to transfer protons to substrates. The relatively recent structure determination for 2a in 2008 has launched a series of pyridinol derived ligands (16a) and metal complexes for both hydrogenation chemistry and other reactions. Parallels in reactivity can readily be drawn between the CpdOH ligands first synthesized by Shvo (Fig. 6) and the 2-pyridinol type ligands that have been used frequently in catalysis. Pyridinol ligands (17a, Fig. 17) can bind to a metal center and upon transfer

Fig. 16 Mono-iron hydrogenase (2a) as modeled by pyridinol metal complexes (16a).

454

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

Fig. 17 2-Pyridone and its tautomers, protonation states, and metal binding modes.

of a proton to substrate the pyridinol ligands become anionic and donate more electron density to the metal center. This parallels the role of the CpdOH ligand in the Shvo catalyst for its ability to change its electron donor properties with ligand protonation state. The ligand 2-pyridinol (17a, pyOH) and its tautomer 2-pyridone (17b) represent the simplest ligands of this type (Fig. 17). The 2-pyridone form is more stable but generally the 2-pyridinol form (or the deprotonated ligand, 17c) is the tautomer that binds to the metal. Once 17c binds to a metal it can be monodentate (O bound (17f) or N bound (17e)) or bidentate (17d) (Fig. 17). The 2-pyridinolate ligand (17c) can also bridge two metals (17g).166 Following a short discussion of metal complexes using 2-pyridinol ligands, we will discuss bidentate, tridentate, and other chelating ligands derived from pyridinol. Coordination complexes of 2-pyridone have been synthesized with most of the middle to late transition metals (groups 6–11) and this topic has been reviewed in 1995,167 but herein our discussion will focus on catalysis with these metal complexes.

1.14.4.2

Ligands containing one pyOH group

Yamaguchi and Fujita reported that monodentate pyridinol derivatives formed highly active iridium catalysts for hydrogenation and dehydrogenation of nitrogen heterocycles (Fig. 18).168,169 Several electron donating and withdrawing substituents were used on the pyridinol ring, and the best results were obtained with the 5-CF3 group (in 18a) giving 100% yield and 100% conversion. It has been suggested that the role of the CF3 group in 18a is that it creates a more acidic OH group to facilitate proton transfer events. Up to five cycles of hydrogenation/dehydrogenation could be performed with near quantitative yield. However, this catalyst was vulnerable to decomposition with the formation of the dimer [Cp IrCl(m-H)]2. Similarly, catalyst 18b has been described as a dehydrogenation catalyst by Rauchfuss and coworkers.170 Under hydrogenation conditions, the main species observed are 18c (the product of H2 addition to 18b) and 18d which represents an inactive catalyst resting state. Wang et al. have studied the mechanism of the reaction with 18a and similar catalysts and have shown that the hydrogenation/dehydrogenation pathways are mediated by reactions on the monomers (e.g. 18a, 18b, and 18c) and that the dimers are inactive resting states. Complex 18a binds H2 as H− on the metal and H+ on the ligand thus allowing for metal-ligand bifunctional catalysis.171 Wang et al. describe this mechanism as involving “ligand rotation-promoted hydrogen transfer pathway” (LRPHT). This mechanism is illustrated in Fig. 18B and is named LRPHT because the ligand switches from being bidentate (18b) to monodentate (18e) for substrate binding, and then H2 release returns the catalyst to a bidentate resting state (18b).172 Thus a coordinatively flexible pyridinol ligand appears to offer some

Fig. 18 (A) Monodenate py-OH iridium complexes for hydrogenation and dehyrogenation. (B) Ligand rotation-promoted hydrogen transfer (LRPHT) pathway for dehydrogenation catalysis.

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

455

Fig. 19 Ruthenium and iridium complexes with phosphine chelated to py-OH.

advantages in catalysis. More recently, several of these metal complexes have been used for the organic synthesis of complex products by catalyzing dehydrogenation which is followed by cyclization in situ.173 Similarly, pyridinol has been linked to phosphine ligands (in Fig. 19 and in other reports174) and used to form iridium and ruthenium complexes; the ruthenium complexes were demonstrated to be catalysts for hydrogenation of ketones and a cooperative mechanism is proposed.175

1.14.4.3 1.14.4.3.1

Chelating ligands containing two pyOH groups Early work: CO2 hydrogenation and other reactions

The problem of metal complex stability and avoiding decomposition and dimerization reactions with pyridinol ligands (as in Fig. 18) is easily solved by using chelating ligands of higher denticity. The work with 4,4ʹ-dihydroxybipyridine (4,4ʹ-dhbp) will be described briefly followed by a more in depth discussion of 6,6ʹ-dihydroxybipyridine (6,6ʹ-dhbp) ligands. Himeda and coworkers have demonstrated that metal complexes of 4,4ʹ-dhbp can be reversibly deprotonated to alter catalyst properties in situ. This allows for proton responsive reactions which results in highly active and reversible catalysts for CO2 hydrogenation.176 These ligands are pH responsive because the form present at low pH (20b, Fig. 20) bears OH groups and is a weaker donor ligand.8,177 In contrast, the deprotonated form, 20a, features O− groups that are stronger electron donors capable of pushing electron density onto the metal center via resonance. The deprotonated 4,4ʹ-dhbp ligand allowed access to the heterolytic MLBC pathway for H2 activation as shown in Fig. 4B, with H+ going to the O− of the ligand and H− on the metal. Furthermore, once the metal hydride is formed, a strong electron donor ligand makes the hydride more hydridic for faster nucleophilic attack on CO2. While Rh, Ir, and Ru all form complexes with 4,4ʹ-dhbp (20c-g) that are hydrogenation catalysts,178 the iridium complexes (e.g. 20d and 20f) are most active for CO2 hydrogenation. This proton responsive ligand has been used to control CO2 hydrogenation in aqueous solution with base present. For example, using catalyst 20f in the presence of base, the deprotonated form (analogous to 20a) is the dominant form in solution for CO2 hydrogenation which generates HCO−2 (formate) (Fig. 3). Base is thus consumed in the reaction and the catalyst switches to its less active form, 20f. At the end of the reaction, when the reaction is vented to release pressure, the catalyst is in a less active form which prevents the back reaction: thereby preventing the catalytic conversion of formate and protons to H2 and CO2.8,177 Most remarkably, [Cp Ir(OH2)(4,4ʹ-dhbp)][SO4] has been used for CO2 hydrogenation to produce formic acid and methanol under mild conditions.179

Fig. 20 Metal complexes of the 4,4ʹ-dhbp ligand.

456

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

Fig. 21 Metal complexes of the 6,6ʹ-dhbp ligand and related ligands with two pyOH groups.

Following this work, Britovsek and coworkers aimed to bring the OH groups closer to the metal center to allow for hydrogen bonding interactions proximate to the metal center to influence catalysis.180 Thus the first 6,6ʹ-dhbp metal complexes were reported in early 2011 using Rh complexes for carbonylation catalysis. The synthesis of metal complexes involved 6,6ʹ-dhbp protection, coordination to the metal, and then deprotection to reveal the OH groups; this procedure was used due to the poor solubility of the 6,6ʹ-dhbp ligand in many commonly used solvents since the protected form had better solubility. The carbonylation of methyl acetate to form acetic acid is reported with moderate yield (35% using catalyst 21a, Fig. 21).180 A follow up study later showed that Rh complexes using bpy ligands displayed similar catalytic rates with five different bpy substituents (including unsubstituted bpy, 6,6ʹ-dhbp, 6,6ʹ-(NH2)2-bpy, etc.). Furthermore, the bpy derivatives actually gave slower catalysis vs simple dimers, e.g. [RhCl(CO)2]2. Thus, this ligand type did not appear to accelerate carbonylation.181 Around the same time, Papish et al. published a paper in late 2011 on ruthenium 6,6ʹ-dhbp complexes (e.g. 21b, Fig. 21) as catalysts for transfer hydrogenation.51 In this paper, protection and deprotection steps for the 6,6ʹ-dhbp ligand were not necessary for metal coordination because the ligand dissolves upon refluxing in DMF, ethanol, water, and other solvents. The dilactam tautomer (analogous to 17b, Fig. 17) is the predominant form of 6,6ʹ-dhbp prior to metal complexation, and this form hydrogen bonds to itself which leads to its poor solubility. However, apparently enough tautomerization occurs in situ (and is likely enhanced by the presense of base)47 to allow for the formation of metal complexes. Complex 21b served as a catalyst for the transfer hydrogenation of ketones in aqueous solution (90:10 of water: methanol), and a bifunctional mechanism was proposed in which the 6,6ʹ-dhbp ligand was the source of H+ and the metal was the source of H− for substrate hydrogenation. In aqueous media, Ru complexes bearing the 6,6ʹ-dhbp ligand gave higher yields of alcohol product than bpy or 6,6ʹ-dimethoxybipyridine ligands suggesting that the OH groups were beneficial. Papish et al. proposed in 2011 that the 6,6ʹ-dhbp ligand could facilitate (de) hydrogenation of other substrates including carbon dioxide by a metal-ligand bifunctional mechanism (Fig. 22).51,182 Recently, derivatives of 21b (e.g. [(Z6-cymene)Ru(6,6ʹ-dhbp)(L)]2+ where L ¼ NH3, OH2) have been used as proton responsive catalysts for ammonia borane solvolysis,183 for the N-methylation of amines,184 and for other reactions. Before describing the large number of papers utilizing typically bidentate pyridinol type ligands (Fig. 21), it is useful to discuss what advantages are offered by 6,6ʹ-dhbp and similar ligands. The geometry of 6,6ʹ-dhbp is well suited for metal ligand bifunctional catalysis (MLBC) with H2 activation occurring as shown in Fig. 22 (22b to 22c to 22d) via heterolytic H2 activation. Furthermore, an electron rich deprotonated ligand assists in H2 activation in 22c (see also Fig. 4). H2 can then be transferred to substrates via a stepwise mechanism with hydride transfer followed by proton transfer. Fig. 22 shows a general schematic for the process, but the individual steps are not shown here (e.g. 22e (no base) and 22f (OH deprotonated by base)).59,185 These steps can be run in the reverse direction for substrate dehydrogenation. In general, the proximity of the pendant base (O− group on 6,6ʹ-dhbp) to the metal hydride helps explain the wide use of this ligand in MLBC of hydrogenation and dehydrogenation.59,185

1.14.4.3.2

Recent work: Hydrogenation, dehydrogenation and related reactions

In 2012, further applications with metal complexes of the 6,6ʹ-dhbp ligand were realized including (de)hydrogenation of challenging substrates. Fujita et al. reported in 2012 that the iridium complex 21c is highly active (92% conversion) and selective for the acceptorless dehydrogenation of benzyl alcohol to form exclusively benzaldehyde.186 A variety of other alcohol substrates were also investigated with similarly high conversion and selectivity. In 2020, Papish, Brewster et al. reinvestigated Fujita’s 2012 report of aqueous benzyl alcohol dehydrogenation with 21c186 and observed only 15% conversion and 98% selectivity (and lower selectivities at lower temperatures) under the reported conditions.187 The observed selectivity was attributed to the reactivity of the aldehyde product including that catalyst 21c leads to further oxidation to form benzoic acid. Benzoic acid is then decarboxylated

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

457

Fig. 22 General qualitative pathways for H2 activation with 6,60 -dhbp metal complexes. The ligand is well suited for heterolytic H2 cleavage via MLBC. H2 transfer to substrates can also occur via a cooperative mechanism. Furthermore, these steps can be run in the reverse direction for substrate dehydrcgenation.

under the reaction conditions to form benzene and CO2. The observed reactivity of both the aldehyde and the carboxylic acid with 21c can be explained in terms of literature precedence involving very similar reactions.46,188,189 Thus, the reaction selectivity differences can be explained, but the exact conditions that may lead to high conversion values for this reaction are unclear. These results187 suggest that while 21c is an exceptional catalyst for several hydrogenation and dehydrogenation reactions, including with CO2 as a substrate as described below, it is not particularly good at benzyl alcohol dehydrogenation in aqueous solution, at least under the conditions reported. Following the initial 2012 report, the Fujita, Li, and other groups have used 21c and similar catalysts for dehydrogenation and hydrogenation of a variety of organic substrates using aqueous, alcoholic, and organic solvents.3,190–195 Kundu and coworkers similarly used [(6,6ʹ-dhbp)RuH(CO)(PPh3)2]Cl as a catalyst for the b-alkylation of secondary alcohols with primary alcohols using a hydrogen-borrowing approach (Fig. 10).196 A similar approach has also been used by Achard and coworkers using catalysts like 21e for b-alkylation via a hydrogen-borrowing approach and for hydrogenation and dehydrogenation reactions that lead to esters and formic acid products.197,198 Similarly, Chen and coworkers have synthesized other ruthenium catalysts with tridentate meridional ligands bearing one or two pyridinol groups.103,199–201 These catalysts have performed ketone transfer hydrogenation and b-alkylation of alcohols via a hydrogen-borrowing approach (Fig. 10). Similarly, iridium catalysts bearing OH groups near the metal (21f, Fig. 21) center are active for the methylation of amines and ketones with methanol as the methyl source.202 Thus, OH groups near the metal center are believed to accelerate catalysis in a wide variety of hydrogenation and dehydrogenation reactions.203–205 Moore and Szymczak have used the new 6,600 -dihydroxyterpyridine ligand in ruthenium complex 21g (Fig. 21) for the transfer hydrogenation of ketones.206,207 Catalyst 21g and related systems have also been used for alcohol oxidation to form carboxylates. However, for benzyl alcohol dehydrogenation NHMes groups were shown to lead to a higher yield of benzoic acid vs OH groups (in 21g) due to steric protection of the coordination environment around Ru.208 The authors proposed that steric bulk prevented the formation of a kinetically inert aquo-bridged dimer.207 Expansion of the chelating ligand from 21g to 21h leads to an ability to form intramolecular hydrogen bonds in 21h (Fig. 21). Catalyst 21h is remarkable for being able to catalyze the hydroboration of ketones and nitriles under mild conditions.209 Catalyst 21h was isolated and fully characterized (including by single crystal X-ray diffraction) in four different protonation states, and catalyst activity was shown to vary with the ligand’s protonation state. However, a later study showed that catalyst similar to 21h but lacking OH groups (but bearing methyl groups in their place) was even more active (by TON) for both CO2 hydrogenation and formic acid dehydrogenation.210 Here, the improvements in TON were attributed to bite angle (21g vs. 21h), ligand charge (−3 when OH groups deprotonated in 21h, vs. −1 with methyl groups analog), and the resulting metal complex charge (which ranged from −3 to neutral).210 Excessive negative charge on the metal complex may negatively impact the H2 binding step in Fig. 4. Overall, it appears that for meridional tridentate ligands on Ru (21g and 21h) provide exceptions to the trend that OH groups lead to the optimum TON values since other substituents are better.

1.14.4.3.3

CO2 hydrogenation and related reactions

Iridium complexes with the 6,6ʹ-dhbp ligand (21c) and related ligands (e.g. 21d)211 are also fast and reversible catalysts for CO2 hydrogenation46,212–214 and electrochemical CO2 reduction.215 The topic of CO2 hydrogenation including with proton responsive ligands has been covered in several review articles.3,23,45,216 These catalysts (21c, 21d, and other related systems) are proton responsive catalysts that are deprotonated under basic conditions.49,214,217 Catalyst 21d is unique for operating at mild temperature and pressures for both CO2 hydrogenation and formic acid dehydrogenation with excellent TON and TOF values (TON as high as 153,000 and 308,000 for CO2 hydrogenation and formic acid decomposition, respectively).211 The closely related

458

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

reaction of methanol production by paraformaldehyde disproportionation was also studied with 21c, catalysts similar to 21d (giving the best results), and related rhodium and ruthenium catalysts.218 Catalyst 21d and related systems are also efficient catalysts for ammonia borane hydrolysis to release H2 gas.219 Catalysts that replace one ring (in 21d) with pyrazole or imidazole derivatives (e.g. 21d0 ) are even more active at CO2 hydrogenation and formic acid dehydrogenation.220–223 The NH groups in 21d0 do not act as proton donors to any significant extent; the impact of the imidazoline moiety in 21d0 simply comes from it being a strong electron donor.222,224 Herein, the tests to determine the function of the NH groups included measuring the pKa values of 8.8–9.0 in aqueous solution by pH dependent UV-Vis (a range is given because a series of closely related compounds were studied).222 When this is compared to the pH for catalysis being 8.3, it appears that partial deprotonation is possible.222 However, a methylated version of the catalyst (NMe vs. NH) performed similarly with nearly identical TOF values. This result ruled out NH deprotonation as influencing catalytic rates. The OH groups are however able to act as proton donors based upon both their pKa values and since their substitution impacts catalytic rates.223 Given the success of catalyst 21d0 with an imidazole derived ring, it seemed logical to test a chelating N-heterocyclic carbene. Catalysts with N-heterocyclic carbene (NHC) rings bound to pyridinol have been reported, and while they do lead to good catalysts for other reactions including dehydrogenation of organic alcohols and N-alkylation of amines with alcohols via a hydrogen-borrowing approach,225,226 they do not offer any special advantage for CO2 hydrogenation.59 In the presence of base for CO2 hydrogenation, the ligand rearranged by cyclometallation and led to less active catalysts. Thus NHC ligands bound to pyridine and pyridine derivatives have found more utility where base is not required or with metals and ligand geometries less vulnerable to cyclometallation.227,228 Examining the role of OH groups in 21c and 21d more closely, catalysts of the type [Cp Ir(n,nʹ-(NH2)2-bpy)(OH2)]2+ (n ¼ 4 or 6) were used for formic acid dehydrogenation.229 Investigating a wide pH range of 2 to 12, the authors were able to show by 1H NMR that NH groups in the 4,4ʹ-(NH2)2-bpy iridium complex are not deprotonated under these conditions. The catalyst with amino groups in the 4,4ʹ positions is more active than the 4,4ʹ-dhbp analogue since the amino groups are electron donating. In contrast, the 6,6ʹ amino groups lead to a less active catalyst vs the 6,6ʹ-dhbp analogue because in this case they are protonated at low pH (forming NH+3 substituents) and lead to an electron withdrawing effect. Effective catalysts using OH groups on a bpy scaffold are not surprising given that Himeda demonstrated that 4,4ʹ-dhbp Ir complexes (e.g. 20d) are highly active for CO2 hydrogenation.8 However, the presence of hydroxy groups near the metal center in 21c leads to significant rate enhancements (vs. 20d with 4,4ʹ-dhbp) that imply that the interaction of OH/O− groups with the substrate serves to lower energy barriers.46 Fujita and Himeda proposed that the O− groups assist by providing a lower energy pathway for H2 heterolysis. Following H2 cleavage, an IrdH group and a pendant OH group are poised to hydrogenate the CO2 substrate.46 This mechanism was further supported by computational studies.185,230 A follow up study by Grotjahn, Papish, Webster and coworkers further clarified the role of the OH groups in the catalysis.59 For CO2 hydrogenation under basic conditions, the OH groups are deprotonated and O− groups can bind Na+ ions (present in solution from the base) which can help activate CO2 (Fig. 23). Both experimental and computational results support the role of both sodium ions and the O− groups. A higher energy pathway was required in the absence of sodium ions computationally. Furthermore, when methoxy groups are used on the catalyst (e.g. with 6,6ʹ-dimethoxybipyridine as the ligand on Ir) the interaction with sodium ions is not as favorable and the computational prediction is that methoxy groups will produce a slower catalyst, as is confirmed experimentally. Transfer of a hydride to CO2 is the rate determining step in each case, and a lower energy pathway is provided by using O− groups (vs. OMe groups) as shown in Fig. 23 (23b shows the highest energy transition state with O− groups). Similar studies for formic acid dehydrogenation to yield CO2 and H2 showed that 6,6ʹ-dimethoxybipyridine Ir complexes analogous to 21c are just as active as 21c with OH groups.59 The presence of formic acid in aqueous solution leads to a lower pH (2) and as such the OH groups in 21c are protonated. The OH and OMe groups leads to mechanisms with similar energy barriers, and herein there is no advantage to the OH groups.59 In many ways, the role of the sodium ions in catalysis under basic conditions herein is similar to the role that potassium ions play with Noyori’s catalyst (5a, Figs. 5, 7, and 8).2 Catalyst 21c also performs a reaction that is the reverse of CO2 hydrogenation. Fujita, Yamaguchi and coworkers have reported that 21c catalyzes the complete dehydrogenation of aqueous methanol to produce CO2 and H2.231 This reaction is proposed to proceed via (1) CH3OH dehydrogenation to formaldehyde, (2) formaldehyde hydration, (3) dehydrogenation to formic acid, and lastly (4) formic acid dehydrogenation to CO2 and H2. Thus, this reaction releases three equivalents of H2 per equivalent of CH3OH.

Fig. 23 Partial mechanism showing key steps in CO2 hydrogenation by 21c in the presence of Na+ ions and base.

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

459

Fig. 24 Dhbp catalysts featuring cobalt, ruthenium, and rhenium metals (24a–24f). These results stand in contrast to catalysis with 24g and 24h.

The best results were obtained with deprotonated 21c and base present producing up to 84% yield of H2 and as high as 10,510 turnovers. A metal-ligand bifunctional mechanism is proposed based on the higher yield of H2 with 21c bearing 6,6ʹ-dhbp vs with analogs bearing 4,4ʹ-dhbp or bpy co-ligands. Similarly, Muckerman, Himeda, Fujita and coworkers investigated 6,6ʹ-dhbp and 4,4ʹ-dhbp Co(III) complexes (e.g. [Cp Co(n,nʹ-dhbp)(OH2)]2+, 24a and 24b, Fig. 24) for activity in CO2 hydrogenation.232 Despite the similarity in structure the iridium analogs, e.g. 21c, the cobalt catalyst featuring 6,6ʹ-dhbp (24b) displayed low activity (TON ¼ 1.3) due to thermal instability. Catalyst 24a was more active and achieved up to 59 TON for basic CO2 hydrogenation. In contrast, the bpy analog achieved only 0.66 TON, illustrating a benefit for inclusion of a proton responsive ligand. However, the cobalt systems have far lower activity vs the iridium catalysts. This reflects a trend in which dhbp complexes of first row transition metals can be less stable vs second and third row metal complexes.233

1.14.4.3.4

Electrocatalytic oxidation and reduction reactions

Ruthenium catalysts using dhbp ligands (similar to 24c and 24d but with Cl or OH2 in place of CH3CN, Fig. 24) were synthesized in 2014 by Paul, Papish, Grotjahn and coworkers and used as potential water oxidation catalysts, but they achieved less than one TON.62 In fact, for water oxidation catalysis, methoxy groups on the bpy ligands were beneficial and led to catalytic turnover (TON as high as 215) but OH groups were detrimental, perhaps due to decomposition events. Catalysts 24c and 24d were studied as catalysts in electrocatalytic CO2 reduction in a 2016 report from Muckerman, Himeda, Fujita and coworkers.234 The analysis of the CV data for 24c was complicated by the formation of an insoluble species upon reduction. However, 24d was more amenable to studies and it was determined to be a non-innocent catalyst due to reactions between the OH/O− groups and CO2 leading to the formation of dOCO−2 substituents on the bpy rings as a decomposition product for 24d. These side reactions served to deactivate the catalyst and limit CO and formate production. Similarly, the rhenium complexes 24e and 24f were investigated as potential electrocatalysts for CO2 reduction.50 However, 24f with the 6,6ʹ-dhbp ligand is not a catalyst but rather undergoes CO loss upon reduction leading to an unstable species. Complex 24e is an electrocatalyst for the reduction of CO2 to CO with 95% Faradaic efficiency. A follow up study examined the 3,3ʹ-dhbp Re analog and showed it to be less active than 24e but more active than 24f.235 However, the use of [4,4ʹ-(tBu)2-bpy)Re(CO)3Cl]236 gives even better results than 24e and thus the OH groups are overall not beneficial for this chemistry. The work on 24a–24f overall shows a detrimental role for OH groups near the metal center specifically for CO2 hydrogenation with cobalt and for electrocatalytic oxidation and reduction reactions with ruthenium and rhenium. The work with 24a and 24b shows the differences in stability of cobalt complexes vs iridium and ruthenium catalysts (e.g. 21b and 21c) and highlights the fact that with first row metals the dhbp complexes can have reduced stability. Similarly, copper complexes of 6,6ʹ-dhbp were shown to demetallate at low and high pH values.233 For 24c–24f, these results also stand in stark contrast to work with other ligands that provide OH groups near the metal center, but in the form of nearby phenol rings. Here, the phenol ring is not coordinated to the metal center, but rather it is simply proximate to the metal and dangling above the CO2 binding site. For example, Bocarsly and coworkers showed that catalyst 24g showed a seven fold current enhancement with (icat/ip)2 ¼ 119 (relative to the bpy analog of

460

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

24g at (icat/ip)2 ¼ 2.6) along with a low overpotential of 440 mV for CO2 reduction to CO.71 The proximate OH group was proposed to assist with CdO bond cleavage by offering an intramolecular hydrogen bond. Similarly, Costentin and Savéant showed that an Fe porphyrin catalyst 24h produced a record setting TOF for electrocatalytic CO2 reduction to CO ((icat/ ip)2 ¼ 3600).22,237 However, these results were not unique to proximate OH groups, and later studies showed even greater rate enhancements with fluorinated alkyl groups and trimethyl ammonium groups decorating the phenyl rings on the porphyrin.9,238,239 Apparently, the use of such polar groups appears to accelerate both proton and electron transfer processes and the trimethylammonium groups improve water solubility.239 Other examples of electrochemical CO2 reduction being accelerated by nearby OH groups from phenols appended to ligands can be found in the literature, including with rhenium catalysts.240–242 The lesson learned here may be that for electrochemical CO2 reduction reactions, it may be helpful to have the protic groups close to the metal center, but not necessarily electronically linked to the metal center.

1.14.4.3.5

Summary of pyridinol based ligands for organometallic catalysis

Pyridinol based ligands described herein form metal complexes that excel at hydrogenation and dehydrogenation reactions. Ruthenium and iridium complexes are featured prominently in this chemistry but first row metals (e.g. cobalt) have also been used. These reactions have included CO2 hydrogenation, formic acid dehydrogenation, and (de)hydrogenation reactions involving organic alcohols, ketones, and carboxylic acids. In some cases, the organic molecules are coupled together using a hydrogen-borrowing approach. Pyridinol based ligands excel here, and appear to be better at CO2 hydrogenation vs the Noyori and Shvo type systems covered in Sections 1.14.2 and 1.14.3. However, moving to oxidation and reduction reactions, the pyridinol based systems are often vulnerable to side reactions. In some cases, the products have been characterized and 6,6ʹ-dhbp can be altered or demetallated in certain reaction conditions.62,233,234 This foreshadows a theme that is repeated in the sections that follow: pyridinol based ligands that are tridentate and facial (Section 1.14.5) or pincer based (Section 1.14.6) are best suited for reactions that do not involve harsh oxidizing or reducing conditions.

1.14.5

Tridentate facial arrangements of hydrogen bond donors or acceptors

Moving away from bidentate and tridentate meridional ligands (Fig. 21), the use of tridentate facial ligands which can donate hydrogen bonds offers some unique advantages,243 which have been discussed in several review articles.5,60,244 One key advantage of the trigonal ligand structure (C3v symmetry) is that it places up to four d electrons in orbitals that are either non-bonding or p in character.245 Thus, metal oxo complexes including those with multiple MdO bonds with mid and late first row metals can be stabilized. Pioneering studies by Borovik and coworkers from 2000 onwards have designed new tripodal ligands with hydrogen bonds near the metal center. The ligand scaffold shown in Fig. 25 (e.g. 25a and 25b) offers urea groups bound to a tren (tren ¼ tris(2-aminoethyl)amine) ligand core. The ligand features which promote the chelation and hydrogen bonding to metal oxo units include: (1) a tridentate and trianionic ligand that can stabilize metals in a high valent state, (2) a rigid scaffold with thermodynamically favored six membered rings formed by the hydrogen bonding interaction (refer to the introduction, Section 1.14.1 and Fig. 5), and (3) the bulky t-butyl groups serve to create a protected cavity.60 This ligand environment supported both MIIIdO (25a) and Mn]O (25b) (n ¼ IV, V) complexes with M ¼ Fe and Mn. The high spin MnV]O complex was the first example isolated and was fully characterized by EPR, resonance Raman and X-ray emission spectroscopy to provide a signature for similar high valent Mn oxo species in enzymes.246,247 While the FeIII complex (25a) initially appears to violate the oxo wall (metal oxo compounds with greater than four d electrons are generally disfavored), in fact the oxo wall concept only applies to octahedral complexes, and the ligand field for trigonal symmetry has higher energies for the p  orbitals. In addition, and most relevant to this chapter, the need for p-bonding is absent in these complexes because they have largely M+dO– character, with the negative charge on the oxygen stabilized by the hydrogen bonds to the ligand.248,249 These iron and manganese (IV) complexes are capable of activating weak CdH and OdH bonds by H atom transfer.250 CdH activation with high valent iron oxo complexes is usually presumed to form high valent iron hydroxide species, but a study of protonation of an FeIV oxo species presents DFT and spectroscopic evidence that the proton can be transferred to the ligand instead.251 CdH activation was also performed with a series of related MnIIIdO complexes in which the hydrogen bond donor strength was modulated by varying an R group which was in conjugation with the NH donor (25f).252 The fastest rate of CdH activation was achieved with R ¼ OMe in 25f because the MnIIIdO starting material is not as stabilized by hydrogen bonding interactions in this case vs with other R groups. Thus the oxo is more basic and reacts more quickly to form 25g. Another noteworthy study involved the isolation of a MnIII-peroxo complex in which the O2 was derived from dioxygen.253 Similarly, Goldberg and coworkers have used a ligand featuring an amide derived hydrogen bond to stabilize a high spin FeIIIOOR species.254,255 Similar studies with cobalt(II) and cobalt(III) complexes like 25c and 25d show that the number of hydrogen bonds present has an impact on the ability of the complex to activate O2. Complex 25d offers three hydrogen bonds, and derivatives were prepared which offer two, one, or zero hydrogen bonds with a similar steric profile. The cobalt(II) complex lacking hydrogen bonds failed to react with O2, and the number of hydrogen bonds offered near the Co(II) center determined whether the complexes reacted slowly (fewer H bonds) or quickly with O2 (more H bonds, e.g. 25c) to generate CoIIIdOH (e.g. 25d).256 Similarly, derivatives with a bridging pyrazolate ligand have been made that bind to two Co(II) ions and offer four hydrogen bonds near the metal centers.257 Manganese and iron complexes with a reversal in polarity have also been synthesized by incorporating hydrogen bond acceptors into a tripodal ligand scaffold denoted as MST (MST ¼ N,N0 ,N00 -[2,20 ,200 -nitrilotris(ethane-2,1-diyl)]

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

461

Fig. 25 Tridentate and facial arrangements of hydrogen bond donors as studied by Borovik, Masuda, Berreau, et al.

Fig. 26 Trigonal arrangements of H bond acceptors synthesized and studied by Borovik et al.

tris(2,4,6-trimethylbenzenesulfonamido) in 26a, Fig. 26). The sulfonamide tripod (in 26a) has been used to synthesize MIIIdOH (M]Fe, Mn) complexes with a group 2 metal ion nearby (MII]Ca, Sr, Ba).258 These complexes (26a) are formed in situ by treating [(MST)MII]− with O2 in the presence of the group 2 metal ion and crown ether present. These reactions were slow in the absence of the group 2 metal ion. Notably, both Ca(II) and Sr(II) lead to significant rate enhancements for O2 activation (as high as 80 fold rate enhancement vs. the reaction lacking the group 2 metal ion). The structure of 26a serves as a structural mimic for the Mn-O-Ca cube edge of the Mn3CaO4 active site within the oxygen evolving complex (OEC) of photosystem II. Spectroscopic studies and cyclic voltammetry studies have suggested that the role of the group 2 ion is to serve as a Lewis acid and remove electron density from the resulting bound hydroxide ligand. However, thus far O2 activation with this ligand scaffold has been stoichiometric rather than catalytic.5,259,260 This scaffold has also proven useful for stabilizing a metal bound NH3 group by hydrogen bonding interactions (26b)261 and for the synthesis of several other bimetallic transition metal complexes.262,263 Ligands analogous to MST, but lacking one arm to offer tridentate coordination and only two hydrogen bond acceptors, have been synthesized and coordinated to FeII and CoII metals.264 The iron(II) complex that results is a catalyst for intramolecular CdH bond amination. In a manner similar to complex 26b, Szymczak and coworkers have incorporated tripodal hydrogen bond acceptors via the O atoms of amide groups into Cu(II) complexes.265 In a structure similar to 26a, the MST ligand has also been used to form cobalt complexes including one formulated as (MST) CoII(m-OH2)CaIIOH2(15-crown-5).266 The oxygen evolving complex (OEC) contains multiple water molecules and an intricate hydrogen bonding network which has been difficult to model in synthetic systems. This aquo bridged Co/Ca structure models the OEC and shows the ability of proton addition/removal to tune the structure and the spectroscopic properties. This species also illustrates the impact of the Ca(II) ion on the redox potentials which are shifted to more positive values by over 200 mV. Furthermore, a related CoIII(m-OH)CaII species is able to activate NH bonds in diphenylhydrazine. Recently, a new ligand scaffold featuring phosphinic amido groups has allowed for the isolation of FeIV]O complexes without hydrogen bonds (26c, Fig. 26) and optionally with a Lewis acid present (26d).267 These structures were notably different from 25b with three hydrogen bonds present. The interaction with the Lewis acid in FeIV(O)—MgII (26d) was shown to weaken the Fe]O bond and the magnesium ion has a similar impact on the structure as three hydrogen bonds in 26d. Furthermore, other ligands have been devised which incorporate both hydrogen bond donors and acceptors and these have stabilized MnndOH complexes with varied oxidation states (n ¼ II, III, IV).268 A series of CoIIdOH complexes with systematically varied hydrogen bond donors and acceptors was also reported.269 This study illustrated how the hydrogen bonds impact the cobalt oxygen bond strength with a greater number of hydrogen bond acceptors leading to a notably short and strong CodO bond.

462

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

Other studies by Borovik et al. have incorporated hydrogen bond donors and acceptors into nickel(II) and zinc(II) complexes. These studies have included tripodal urea based donors bound to Ni(II)dOH (similar to 25d) and tripodal sulfonamide donors bound to Ni(II)dOH2 (similar to 26a).270 Tridentate ligands that donate two hydrogen bonds via urea groups have been used to support square planar Ni(II) complexes as well; these Ni(II) complexes incorporate two internal hydrogen bonds.271 Tripodal Zn(II) and Fe(II) complexes with three internal H bonds to the NH groups of amino-oxazoline based ligands have also been synthesized and studied.78 Gilbertson and coworkers have also studied Fe(II)272–275 and Zn(II)79 complexes with internal hydrogen bonds that in some cases accelerate catalysis. Similarly, Masuda, Berreau, Goldberg, Karlin, and Mareque-Rivas have studied late transition metal complexes that incorporate hydrogen bonds in a trigonal bipyramidal coordination environment.276–284 Masuda’s isolation of 25e was the first example of a structurally and spectroscopically characterized CuIIdOOH complex as a model for copper oxygenase enzymes.284,285 Here the reactive hydroperoxo unit is stabilized by two hydrogen bonds provided by the ligand. In a further example, Berreau et al. have studied 25h which is a stoichiometric reagent for hemithioacetal isomerization as a model for the glyoxalase I enzyme.283 When the basic amide group on 25h is protonated, it is unable to initiate this reaction, thus demonstrating proton dependent reactivity. This was rationalized in terms of the basic amide group being essential for hemithioacetal deprotonation. Earlier work by Berreau and coworkers also established that hydrogen bonding interactions serve to stabilize zinc alkoxide species that are relevant to liver alcohol dehydrogenase enzymes.278 Hydrogen bonds have also been incorporated into neutral tripodal ligands. These ligands contrast with Borovik’s systems which often feature trianionic ligands as in 25a-d (Fig. 25). Baldwin and coworkers reported an early example of a tripodal ligand which offers hydrogen bonds near the metal center in 27a (Fig. 27).286,287 This Ni(II) complex forms biomimetic complexes with nitrate and can be studied in different ligand protonation states. However, a lack of steric bulk made this complex vulnerable to the formation of dimers. Rauchfuss and coworkers have designed complex 27b (Fig. 27) which undergoes reversible deprotonation to create a more electron rich ligand and shift the CO stretch by 26 cm−1.288 A follow up study showed that hydrogen bonding amino and pivaloyl groups in the secondary coordination sphere of copper complexes were of neutral impact (overpotential unchanged) for the oxygen reduction reaction.289 Szymczak and coworkers have designed tripodal ligands based upon 2-hydroxy-pyridine donor groups (e.g. 27c-f, Fig. 27). Parallels to the chemistry in Section 1.14.4 can be readily drawn including that pyridinol changes both charge and donor properties upon deprotonation (see Figs. 17, 20–22). Complexes 27c and 27d show hydrogen bonding between the chloride and the OH groups and exhibit a significant shortening (0.3 A˚ ) of the CudCl bond upon oxidation of 27c to form the Cu(II) complex 27d.290 In the case of the fluoride complex 27e, reduction of Cu(II) to Cu(I) results in cleavage of the CudF bond, but the fluoride is retained by hydrogen bonding interactions to the ligand.291 The use of hydrogen bonds to bind guest molecules is attractive.292 Complex 27f has also been used as an effective reagent for nitrite reduction to produce nitric oxide gas in a reaction that mimics the nitrite reductase enzyme chemistry.293 Switching to NH hydrogen bond donors in 27g resulted in a Cu(I) complex that was capable of binding O2 to form a peroxo complex.294 Herein, the redox potential could be varied by changing the hydrogen bond strength via changes to the R groups. Similarly, NH hydrogen bonds on a tripodal scaffold (27h) were also used to stabilize Ni(I) halide complexes.295 This same ligand scaffold was used with Zn(II) (in 27i in which X is absent to provide a vacant site) to promote the binding and release of O2.296

Fig. 27 Trigonal ligands which offer hydrogen bonds near the metal center.

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

463

Similarly, Fout and coworkers have used tripodal NH donors (in 27j, Fig. 27) for catalytic chemistry including nitrate reduction to NO and perchlorate reduction to chloride and water.297,298 This ligand can act as either a hydrogen bond donor or a hydrogen bond acceptor depending on the protonation state of the NH groups.299,300 This ligand supports a stable ring size for the hydrogen bonding interaction, as a seven member ring results from the ligand being either a hydrogen bond donor (O. . .NH in 27j) or acceptor (OH. . .N in 27k). Furthermore, the NH group is conjugated with the pyrrole N that coordinates to iron via the double bonds. Thus, upon NH deprotonation, more electron density is donated to iron. This scaffold is also electronically tunable, as both the Cy (Cy ¼ cyclohexyl) group and the substituents on the pyrrole ring can be modified to shift the redox potential by 400 mV.301 With nitrite as a ligand, they were able to isolate an O-bound nitrite complex with Zn(II), and here the O-bound binding mode is supported by an NH hydrogen bond to the O of nitrite.298 Using an Fe(II) triflate starting material with the same tripodal ligand, they were able to achieve as high as 3 TON for nitrate and perchlorate reduction reactions with diphenylhydrazine as the source of electrons and protons.297 Stoichiometric nitrite reduction and oxygen activation have also been reported.298,302 The iron containing products obtained included both 27j (with FeIIIdO) and, in the case of perchlorate reduction, a product containing FeIIdCl bound to the tripodal ligand. Complex 27j readily regenerates the active catalyst by forming water with a source of protons and electrons. Hydrogen bonds appear to facilitate this reaction. The catalytic efficiency was improved in a follow up paper by sequestering the chloride formed by binding it to a Zn complex using the same tripodal ligand and by switching to a THF solvent system; thus a TON of 94 could be obtained.303 However, in this reaction the Fe complex is a catalyst and the Zn complex is a stoichiometric reagent which uses an expensive ligand. Furthermore, Fout and coworkers have recently reported a new ligand which offers two hydrogen bonds and five nitrogen donors to iron, as shown in 27k (Fig. 27).304 This ligand has supported Fe(II) and Fe(III) complexes with OH and OH2 ligands hydrogen bonded to the NH groups. Two different protonation states could be isolated for the Fe(III) hydroxo complexes (as shown in 27k, and also the protonated analog with one more NH group present). Both of the Fe(III) hydroxo complexes are capable of abstracting hydrogen atoms from weak CdH bonds with concomitant OH2 formation. However, only the protonated Fe(III)dOH complex was capable of transferring •OH to the triphenylmethyl radical via a radical rebound mechanism. These mechanisms are relevant to the activity of iron containing non-heme enzymes that activate O2 for substrate oxidation. Moving beyond iron, this same tripodal ligand also forms a MnIIIdO complex analogous to 27j, with O2 as the source of oxygen.305 Manganese(II) aqua and hydroxo complexes have also been isolated and characterized.305 Fout and coworkers have also designed other ligands which offer two hydrogen bonds in proximity to late first row transition metals (e.g. Fe, Cu).306 Kuwata et al. reported another system with a tripodal arrangement of NH donors as shown in 27l.307 Again, a lack of steric bulk (similar to the Baldwin example in 27a)286 meant that dimer formation was possible, and a dimer, [(LRuCl)2]+, was used as the starting material for catalysis. The dimer was proposed to dissociate and generate 27l during catalytic hydrazine disproportionation. Up to 9 turnovers (per mole of Ru) was reported with this catalyst.307 Summarizing Fig. 27, every catalyst or reagent therein utilized NH or OH hydrogen bonding groups that were in conjugation with the bound metal via p bonds. Thus, deprotonation served to electronically tune the metal center. Hydrogen bonds, when present, typically led to stable six or seven member rings (with 27f providing an exception with a 10 membered ring due to a dangling fluoride).291 It appears that the optimum system is sterically protected to prevent side reactions. For these reasons, 27j worked very well and was catalytically active for redox reactions described above.297–300 Some benefit may be anticipated by reducing the steric bulk enough to permit the approach of bulkier substrates, for example with 27k.304 However, in practice 27k did not achieve catalysis yet, but it does abstract weak CdH bonds. The design of new catalyst molecules accordingly must balance a need for steric protection to prevent side reactions against a need for access to the metal center for catalysis. Alternatively, hydrogen bonding interactions that are oriented away from the metal binding pocket can also tune the chemistry at the metal center. The tris(1,2,4-triazolyl)borate (Ttz) anion (28a, Fig. 28) typically binds to a metal center via the endodentate nitrogen atoms which leaves the exodentate nitrogen atoms free to bind to protons or additional metal ions. This ligand was first reported in 1966 by Trofimenko and was later studied by Janiak, Papish, and others.308–310 This topic has been extensively reviewed elsewhere and is covered briefly here.70 Complex 28b was studied by Papish et al. and provided an opportunity to monitor both the spectroscopic properties and the catalytic ability of a (TtzR,Rʹ)CuCO complex in the presence and absence of added protons.69 The addition of up to two equivalents of protons gradually shifted the CO stretch in the IR spectrum from 2086 to 2117 cm−1. This change in the IR spectrum was attributed to binding up to two protons to the exodentate nitrogen atoms on the TtztBu,Me ligand. This electronic influence was used to generate a more electrophilic catalyst and enhance the ability of a derivative of 28b, namely the catalyst [(TtztBu,Me)Cu(NCCH3)], to perform a CdH activation reaction between cyclohexane and ethyl diazoacetate. Here the yield increased from 10% to 40% for the CdH activation product when the catalyst was protonated with two equivalents of H+. These results are similar to the usage of fluorinated or halogenated catalysts for CdH activation,311,312 but here the electronic adjustment is done in situ with a proton source. Other publications have reported on the rich chemistry of Ttz ligands with late first row transition metals,74,75,313–318 including for CO2 activation77 and nitrite reduction.319,320 Other applications of Ttz ligands have included the use of (Ttz)Ru complexes (e.g. 28c) for base free transfer hydrogenation catalysis to convert ketones to alcohols.321 Here, steric bulk appears to be important for the reaction to proceed and to allow the free triazole ring (due to Ttz binding k2) to serve as a pendant bases. Machan and coworkers have studied K[(Ttz)M(CO)3] and (Ttz)M(CO)2(NO) complexes (M]Mo, W) and have demonstrated that they are tuned by protonation at the exodentate nitrogen atoms resulting in a shift in the redox potential towards more positive values by 200 mV.73 Cundari et al. published a computational study demonstrating that for Ttz transition metal complexes whether proton coupled electron transfer occurs sequentially or via a concerted pathway is dependent on the particular metal used.322 They argue that these trends can be useful for the design of electrocatalysts to activate small molecules.

464

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

Fig. 28 Ttz ligands that are tuned by exodentate hydrogen bonding interactions (28b) or can act as pendant bases in catalysis (28c). Furthermore, other ligands in 28e, 28f, and 28g also offer pendant basic groups.

Similar to the Cu complexes of the Ttz ligands shown above, several other ligands have featured pendant basic groups that are believed to be important for catalysis.323 Periana and coworkers have designed platinum complexes of bispyrimidine (e.g. 28d) for CH activation of methane.68,324,325 Here, protonation of the remote nitrogen atoms helps produce a more electrophilic platinum center for CdH activation. Similarly, Dubois (28e),326,327 Mayer (28f),328 Wiedner,329–332 Linehan, and others333 have studied systems featuring metals with nearby pendant basic amine groups capable of facilitating hydrogen oxidation and the reverse reaction along with other proton shuttling reactions. A detailed coverage of these topics is beyond the scope of this chapter, but the reader can refer to several review articles.334,335 In summary, this section has focused on the activation of small molecules including O2 activation, NOx reduction, and CdH activation. Significant effort has been invested in designing tripodal metal complexes which offer hydrogen bond donors or acceptors near the metal center. Often, these metal complexes are biomimetic and resemble enzyme active sites. These studies have included both catalytic transformation and stoichiometric activation of small molecules with manganese, iron, cobalt, and copper complexes being frequently used. Nickel and ruthenium metals have been used less frequently in this section. These tripodal protic ligands offer advantages for stabilizing unusual intermediates and facilitating proton transfer in reactions. Likewise, hydrogen bond acceptor ligands serve to reverse the polarity of the pocket above the metal and are well suited to binding small molecules like NH3 and for binding Lewis acids.

1.14.6

Proton responsive pincer ligands

Pincer ligands offer three donor groups to a metal center with meridional binding in octahedral symmetry. In contrast to the tripodal ligands with C3v symmetry and facial binding to a metal center in the prior section, pincer ligands and their metal complexes tend to be used for different types of reactions. Hydrazine disproportionation, carbon dioxide reduction, and hydrogenation reactions are featured prominently in this section. The metals typically used include iron, ruthenium, nickel, and others. In this section, we discuss pincer ligands that contain protic pyrazole rings and protic NHC rings. Therefore, other closely related protic ligands, namely those containing protic NHC or pyrazole rings in bidentate ligands, are briefly mentioned due to the similarity in structure to protic NHC containing pincers. Pincer ligands were already discussed in the context of ligands that offer two pyridinol derived OH groups near the metal center as in ruthenium complexes 21g and 21h (Fig. 21). However, in this section we focus on ligands for which the protic OH group is further from the metal or the protic group is an NH substituent. This includes ligands bearing pyrazole NH groups near the metal center bound to ruthenium, iron, and other first row transition metals. This also includes CNC and NNN pincer ligands bound to copper, nickel, and ruthenium with OH groups for proton responsive catalysis. Kuwata, Ikariya, and coworkers described the Fe(II) complex 29a (Fig. 29) which features an NNN pincer ligand bearing two NH groups from pyrazole rings near the metal center.336 This catalyst (20 mol% relative to substrate) was capable of converting hydrazine to ammonia and dinitrogen gas by a disproportionation reaction (>98% yield). A mechanism was proposed in which the NH group was a proton source for forming NH3, and the resulting pincer incorporating pyrazolate also served as an internal basic site (attacking R of NH2NR2) to facilitate N2 formation. Following this initial report, Lord, Caulton, and coworkers further

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

465

Fig. 29 Pincer ligands and related metal complexes offering proton responsive NH groups near the metal center.

studied a closely related (NNN)Fe(II)Cl2 complex with the same pincer ligand.337 This complex showed irreversible redox events with the protic NH groups present, but methylation of these groups led to reversible redox behavior. Furthermore, a one electron reduction with KC8 was performed on the potassium salt of the free ligand and EPR was used to show that the localization of electron density was primarily on the pyridine ring in the resulting ligand based anion radical. Studies by Caulton and coworkers on the related [CrII(NNN)2] complexes using the same pincer ligand as in 29a show that these chromium complexes are able to promote stoichiometric hydrazine disproportionation.72,338 Related Ru pincer complexes have also been studied for formic acid dehydrogenation, ketone hydrogenation, and small molecule activation.339–342 Furthermore, Caulton and coworkers have synthesized Co(II) and Co(III) complexes with protic bis(pyrazole)pyridine (NNN) pincers.343 One example of a Co(III) complex is shown in complex 29b (Fig. 29). A follow up paper showed that these Co complexes can form multimetallic aggregates via hydrogen bonding interactions as well as pyrazolate coordinating to two metals via bridging interactions.344 Related protic PNN ligands bearing pyrazole NH groups have been reported which form Co(I/II) complexes.345 One Co(II) complex performs stoichiometric nitrite reduction to generate Co(III) products including one NO complex.346 Computational studies have supported the participation of the protic NH group in nitrite reduction. Specifically, proton transfer from the pyrazole NH to the NO2 ligand is feasible and is postulated as a required step for NdO bond cleavage. This ligand geometry supports hydrogen bond formation between NH and the O of nitrite, with a computed H. . .O distance of 1.948 A˚ .346 Further studies of protic pincer complexes of ruthenium(II) have included CNN ligands based upon protic NHC-pyridinepyrazole chelates.347 Despite the protic NHC ligand (in 29c), the pyrazole ring is more acidic and deprotonation of the pyrazole group increases electron donation towards Ru. Protic NCN ligands and their Ru(II) complexes were also described with the carbon donor from a central NHC ring and the N donors being protic pyrazole groups.348 Similarly, NCN pincers based upon a central phenyl carbon donor and pyrazole rings coordinated to Ir(III) and Ru(II) have also been described and some of these complexes have been used as catalysts for transfer hydrogenation of ketones.349 Likewise, bidentate CN and NN ligands bearing protic pyrazole rings have formed effective iridium and ruthenium catalysts for intramolecular hydroamination and amine dehydrogenation, respectively.336,350–354 With NH groups near the metal center, these catalysts resemble Noyori’s catalysts and closely related systems (Fig. 7). Here, NH proximity to the metal center combined with an appropriate pKa value (for the pyrazole NH groups) explains the advantages of 29c and related systems. Protic NHC ligands have been frequently used in catalysis both as catalytic intermediates and as the starting metal catalyst. Ellman, Bergman, and coworkers used N-heterocycles including imidazole derivatives as substrates in rhodium catalyzed CdH functionalization, and these reactions were shown to involve protic NHC complexes as intermediates.355–362 Imidazole containing substrates are able to bind to metals and form the protic NHC tautomer as a key intermediate during the CdH functionalization reaction. Similarly, iridium, ruthenium, rhodium, and osmium form protic NHC complexes (Fig. 29, compounds 29d, 29e, and 29f) that have been used as catalysts.350,363–372 The pioneering research groups in this area have included Ikariya, Grotjahn, Schanz, Hahn, and others. However, generally speaking traditional NHC metal complexes with N-alkyl rather than NdH substituents are more frequently used in catalysis, perhaps due to the stability of NHC rings with greater steric bulk which inhibits dimer formation. A full coverage of protic NHC ligands in organometallic catalysis is beyond the scope of this chapter, but the reader is directed to several recent reviews.373–376 In some systems, the pincer ligand bears a protic group further from the metal center. Early work has included terpy metal complexes that feature an OH and other groups in the para position on the central pyridine ring.377–379 Substitution at this position is known to have a strong inductive influence on the metal center. Divalent cobalt, iron, and ruthenium complexes have been reported and studied by spectroscopic and crystallographic methods.377–379 However, they have not been used for catalysis or for proton responsive properties, and so they are only mentioned briefly here.

466

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

Fig. 30 Pincer ligands offering proton responsive OH groups further from the metal center.

Webster, Delcamp, Papish, et al. reported a CNC pincer ligand in 2018 which forms proton responsive Ni(II) complexes (e.g. 30a, Fig. 30).380 The O− group can be protonated to form the OH group with corresponding spectral changes to suggest that the OH bearing cationic complex has less electron density at Ni(II). Catalyst 30a is an active catalyst (TON ¼ 10.6) for photocatalytic CO2 reduction to form CO with visible light, photosensitizer, and sacrificial electron and proton sources present. In contrast, protonation of 30a to give the OH bearing cationic Ni complex leads to a loss of activity. Similarly, the analog with H on the pyridine ring is also inactive (0.9 and 0.1 TON, with OH and H on the pyridine ring, respectively).380,381 Thus, deprotonation switches this catalyst on, and protonation switches it off due to the formation of an inactive complex. This study suggested that the rationale for these results included that photocatalytic reduction of 30a leads to an intermediate that is thermodynamically better positioned to reduce CO2 with a more negative reduction potential, relative to the analogous intermediates with different substituents on the pyridine ring (OH, H). Furthermore, computations showed that the intermediates on the pathway for CO2 reduction showed a greater interaction and better binding between CO2 and the reduced nickel species generated from 30a, vs from other analogs. Despite the promising results with the O− substituent on pyridine relative to OH or H substituents, unpublished results suggest that the analog of 30a with an OMe group on the pyridine ring is much more active for photocatalytic CO2 reduction. These results fit with the studies using Ru for photocatalytic CO2 reduction (see 30b) and the studies described in Section 1.14.4 using Ru and Re dhbp complexes for electrocatalytic CO2 reduction (e.g. 24c-f). Similarly, Papish, Delcamp, and Webster reported ruthenium catalyst 30b with an OH substituent. This catalyst had very low activity for photocatalytic CO2 reduction and this was attributed to in situ decomposition.63 Often, photocatalytic CO2 reduction reactions can be very sensitive to substituent effects, and substituents that resemble sacrificial donors (e.g. NMe2, OH, etc.) often do not perform well. For Ru(II) CNC pincer catalysts of type 30b, the most effective pyridine substituent for photocatalytic CO2 reduction was typically a para OMe group for good p electron donation without as facile side reactions.63,382–384 This trend is similar to what has been observed for Ni(II) complexes of these same pincer ligands (30a, see discussion above).380 Frequently, there is a contrast between the substituent trends for hydrogenation reactions vs electrocatalytic and photocatalytic redox reactions. Repeatedly, we have observed that dihydroxybipyridine (dhbp) complexes excel at hydrogenation reactions, but they are not well suited for electrocatalytic CO2 reduction (Section 1.14.4). Again, this trend repeats and we see that pincer complexes like 30b are better suited for thermal hydrogenation reactions (below), but 30b is a poor catalyst for photocatalytic CO2 reduction. It is clear that OMe groups are a useful substituent for electrocatalytic and photocatalytic redox reactions due to electron donation without a tendency towards side reactions, but OH groups tend to lead to frequent side reactions and decomposition in the case of photocatalytic reactions. The utility of OH/O− groups in thermal hydrogenation catalysis is tied to the ability of strong donor ligands to promote heterolytic H2 activation and to increase the hydricity of the resulting hydride ligand (Fig. 4 above). Switching gears and investigating thermal hydrogenation reactions with these protic pincer ligands, the advantages of OH substituents become more clear. Papish, Vannucci, and coworkers used catalyst 30b which contains a CNC pincer ligand which can allow for different protonation states via an OH/O− substituent on the pyridine ring. This Ru(II) catalyst was highly active for the hydrodeoxygenation (HDO) reaction of vanillyl alcohol to creosol with H2 gas.61 This reaction was enhanced by the presence of base, and up to 10,000 TON were achieved. The base is proposed to deprotonate the catalyst 30b, which is consistent with the observed spectral changes including the appearance of a C]O stretch in the IR, upfield shifts in the 1H NMR for the pyridine ring CH resonances, and crystallographic studies. Other electron rich substituents are also beneficial for the HDO reaction (e.g. OMe, NMe2 on the pyridine ring), but 30b with in the presence of base was the most active in terms of TON and TOF presumably due to the greater electron donating ability of the O− group. This increased electron donation was proposed to enhance trans labilization of a ligand (CH3CN) to create a free site for H2 binding and then subsequent heterolytic H2 activation (Fig. 4). Zhang, Peng, and coworkers designed ruthenium catalyst 30c to offer multiple protic sites including OH and NH groups.385 This catalyst was used for CO2 hydrogenation in THF/water under basic conditions, and a comparison to an analog lacking the OH group showed improved TON values with the OH group (up to 1370 for 30c vs. 1140 with 4-H on the pyridine). However, they also used a catalyst lacking OH groups but with PPh3 ligands (30d), and this catalyst performed even better (TON ¼ 1530). This shows the beneficial effect of strong donor phosphine ligands. They did not test a catalyst bearing both phosphine and OH groups, to see if further TON enhancements were possible. In this work, OH groups improved TON values, but not by a large amount.

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

467

Green and coworkers have synthesized copper(II) complexes of type 30e with the pyridine substituent being OH (shown), OBn, H, Cl, or NO2.386 They have used NMR, single crystal X-ray diffraction, cyclic voltammetry, and DFT studies to determine the impact of electron donating vs withdrawing substituents on the structures. Herein, OH and OBn were the most electron donating, but the OH groups were not deprotonated to access the even more electron rich O− substituent. Computational studies showed that the electron rich OH and OBn substituents led to an increased (relative to H analog) HOMO-LUMO gap for the frontier orbitals, and the electron withdrawing NO2 group has the smallest the HOMO-LUMO gap. This trend is also reflected in the Cu(II/I) reduction potential, which is shifted more negative by 194 mV for 30e vs the H analog. For future studies, this new NNN ligand and its metal complexes may have applications in catalysis with hydrogenation being an attractive direction with middle transition metals (Ru, Ir, Fe) and nitrite/nitrate reduction or other redox reactions being attractive with copper. Other pincer ligands that are capable of donating H+ include PNP, CNC, PNN, PNNP, and other ligands. Many of these ligands feature an acidic CH group in the ligand backbone.387–394 Some of these ligands alternatively donate H+ from a metal bound NH group.395 These complexes have been used for metal-ligand cooperative (de)hydrogenation catalysis with metals such as iridium, ruthenium, and iron. Morris, Milstein, Periana, and their respective groups have reported novel ligands and their iron and ruthenium complexes in which NH groups or acidic CH groups can lead to dearomatization/rearomatization to accelerate hydrogenation catalysis and other reactions.393,396–398 Similarly, the Broere group has synthesized a new proton-responsive “expanded” pincer ligand with a PNNP core that can host two copper centers with three reversibly obtainable protonation states, concomitant with the partial and full dearomatization of the naphthyridine core. The fully dearomatized-ligand Cu complex can also stoichiometrically activate H2 cooperatively to form a dimer with a tetranuclear Cu4H2 core.399 Many of the catalysts listed above are highly active and offer interesting examples of ligands accelerating catalysis by the participation of protic groups. Further detailed discussion is beyond the scope of this chapter, but review articles on this topic are available.1,3,64 In summary, protic pincer ligands have been used to form metal complexes with iron, ruthenium, nickel, and cobalt. These pincer ligands have incorporated protic pyrazole, NHC, and pyridinol rings and have been most useful for thermal reactions including small molecule activation, hydrazine disproportionation, and hydrogenation reactions. In general, the protic groups in the form of pyridinol substituents have been detrimental towards the activity of catalysts used for photocatalytic CO2 reduction with better results being observed for OMe substituents which allow for electron donation without a tendency towards side reactions.400

1.14.7

Perspective

Protic ligands have found utility in many different aspects of organometallic catalysis. Several highly active “privileged” catalysts have incorporated NH or OH groups near the metal center. Bidentate and tridentate meridional (pincer) ligands have been incorporated into highly active ruthenium, iridium, and iron catalysts for hydrogenation and other reactions. The well known systems to use NH groups near the metal center have included Noyori and Ikariya catalysts which offer fast rates and high enantioselectivity for hydrogenation reactions. Various pincer ligands have also used NH groups from pyrazole or NHC rings near the metal center. The first systems to incorporate OH groups near the metal center have included the Shvo system and related iron derivatives incorporating a CpOH ring. Such Shvo type systems ultimately inspired the development of pyridinol based ligands, including dihydroxybipyridine and pincers, which led to highly active catalysts for both CO2 hydrogenation and for the hydrogenation of organic substrates. Often these systems with chelating pyridinol based or protic pincers coordinated to ruthenium, iridium, and other metals excelled at thermal reactions but failed to produce stable catalysts when used for photocatalytic and electrocatalytic redox reactions. Finally, hydrogen bond donors (e.g. OH, NH) and hydrogen bond acceptor groups have also been incorporated in tridentate and facial scaffolds using first row metals and ruthenium. These trigonal metal complexes have typically excelled at small molecule activation (e.g. O2, NH3) using manganese, iron, and copper in environments that resemble enzyme active sites. The use of hydrogen bond acceptors has also allowed the binding of Lewis acids near the transition metal in mimics of photosystem II. Thus, protic ligands have proven to be advantageous for organometallic and other transformations using both earth abundant and precious metals in geometries inspired by the best natural and man-made catalysts. Overall, the factors that influenced the success of proton responsive and hydrogen bonding ligands have frequently included the acidity of the pendant protons, electronic conjugation between the protic groups and the metal center, and ligand rigidity to hold the protic groups at an appropriate distance. Furthermore, efforts to avoid decomposition products by steric protection or selection of metals and ligands with slower exchange rates has been a key factor in the longevity of any given catalyst. These factors allow access to a heterolytic metal ligand bifunctional pathway for H2 activation (Fig. 4), and these factors also allow the activation of other small molecules. Depending on the selection of ligands, metals, and the appropriate reaction to catalyze, the inclusion of protic and hydrogen bonding ligands can be beneficial or detrimental to catalysis. This chapter describes many strategies to select systems in which protic and hydrogen bonding groups lead to catalytic rate accelerations.

Acknowledgments ETP thanks NSF (CHE-1800214 and CHE-2102416) for financial support. We thank Olaitan Oladipupo for assistance with literature searching.

468

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

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. 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. 54. 55. 56. 57. 58. 59. 60. 61. 62. 63. 64. 65. 66.

Khusnutdinova, J. R.; Milstein, D. Angew. Chem. 2015, 127, 12406–12445. Dub, P. A.; Gordon, J. C. Dalton Trans. 2016, 45, 6756–6781. Shimbayashi, T.; Fujita, K.-i. Catalysts 2020, 10, 635. Wodrich, M. D.; Hu, X. Nature Rev. Chem. 2017, 2, 0099. Cook, S. A.; Borovik, A. S. Acc. Chem. Res. 2015, 48, 2407–2414. Rakowski DuBois, M.; DuBois, D. L. Chem. Soc. Rev. 2009, 38, 62–72. Shook, R. L.; Borovik, A. S. Chem. Commun. 2008, 6095. Himeda, Y. Eur. J. Inorg. Chem. 2007, 2007, 3927–3941. Costentin, C.; Robert, M.; Savéant, J.-M. Acc. Chem. Res. 2015, 48, 2996–3006. Crabtree, R. H. The Organometallic Chemistry of the Transition Metals; Wiley, 2014. Xu, T.; Yin, C.-J. M.; Wodrich, M. D.; Mazza, S.; Schultz, K. M.; Scopelliti, R.; Hu, X. J. Am. Chem. Soc. 2016, 138, 3270–3273. Nicolet, Y.; de Lacey, A. L.; Vernède, X.; Fernandez, V. M.; Hatchikian, E. C.; Fontecilla-Camps, J. C. J. Am. Chem. Soc. 2001, 123, 1596–1601. Tye, J. W.; Lee, J.; Wang, H.-W.; Mejia-Rodriguez, R.; Reibenspies, J. H.; Hall, M. B.; Darensbourg, M. Y. Inorg. Chem. 2005, 44, 5550–5552. Murray, K. A.; Wodrich, M. D.; Hu, X.; Corminboeuf, C. Chemistry 2015, 21, 3987–3996. Carlson, M. R.; Gilbert-Wilson, R.; Gray, D. R.; Mitra, J.; Rauchfuss, T. B.; Richers, C. P. Eur. J. Inorg. Chem. 2017, 2017, 3169–3173. Ghosh, S.; Hogarth, G.; Hollingsworth, N.; Holt, K. B.; Richards, I.; Richmond, M. G.; Sanchez, B. E.; Unwin, D. Dalton Trans. 2013, 42, 6718–6775. Barton, B. E.; Rauchfuss, T. B. Inorg. Chem. 2008, 47, 2261–2263. Liu, T.; Darensbourg, M. Y. J. Am. Chem. Soc. 2007, 129, 7008–7009. Liu, X.; Ibrahim, S.; Tard, C.; Pickett, C. Coord. Chem. Rev. 2005, 249, 1641–1652. Moore, C. M.; Dahl, E. W.; Szymczak, N. K. Curr. Opin. Chem. Biol. 2015, 25, 9–17. Rakowski Dubois, M.; Dubois, D. L. Acc. Chem. Res. 2009, 42, 1974–1982. Costentin, C.; Drouet, S.; Robert, M.; Savéant, J.-M. Science 2012, 338, 90–94. Wang, W.-H.; Himeda, Y.; Muckerman, J. T.; Manbeck, G. F.; Fujita, E. Chem. Rev. 2015, 115, 12936–12973. Olah, G. A.; Goeppert, A.; Prakash, G. K. S. Beyond Oil and Gas: The Methanol Economy, 2nd ed.; Wiley-VCH: Darmstadt, Germany, 2009. Elgrishi, N.; Chambers, M. B.; Wang, X.; Fontecave, M. Chem. Soc. Rev. 2017, 46, 761–796. Liu, W.-C.; Baek, J.; Somorjai, G. A. Top. Catal. 2018, 61, 530–541. Goeppert, A.; Czaun, M.; Jones, J.-P.; Prakash, G. K. S.; Olah, G. A. Chem. Soc. Rev. 2014, 43, 7995–8048. Kristjánsdóttir, S. S.; Norton, J. R. Acidity of Hydrido Transition Metal Complexes in Solution, Chapter 9; VCH: New York, 1992. Kubas, G. J. Chem. Rev. 2007, 107, 4152–4205. Papish, E. T.; Magee, M. P.; Norton, J. R. Recent Advances in Hydride Chemistry; Elsevier, 2001; pp 39–74. Crabtree, R. H. Acc. Chem. Res. 1990, 23, 95–101. Morris, R. H.; Sawyer, J. F.; Shiralian, M.; Zubkowski, J. J. Am. Chem. Soc. 1985, 107, 5581–5582. Morris, R. H. Coord. Chem. Rev. 2008, 252, 2381–2394. Chirik, P. J. Acc. Chem. Res. 2015, 48, 1687–1695. Friedfeld, M. R.; Shevlin, M.; Margulieux, G. W.; Campeau, L.-C.; Chirik, P. J. J. Am. Chem. Soc. 2016, 138, 3314–3324. Chirik, P. J.; Wieghardt, K. Science 2010, 327, 794–795. Yu, R. P.; Darmon, J. M.; Milsmann, C.; Margulieux, G. W.; Stieber, S. C. E.; DeBeer, S.; Chirik, P. J. J. Am. Chem. Soc. 2013, 135, 13168–13184. Pritchard, J.; Filonenko, G. A.; van Putten, R.; Hensen, E. J. M.; Pidko, E. A. Chem. Soc. Rev. 2015, 44, 3808–3833. Housecroft, C. E.; Sharpe, A. G. Inorganic Chemistry, 4th ed.; Prentice Hall: UK, 2012. Uematsu, N.; Fujii, A.; Hashiguchi, S.; Ikariya, T.; Noyori, R. J. Am. Chem. Soc. 1996, 118, 4916–4917. Noyori, R.; Sandoval, C. A.; Muñiz, K.; Ohkuma, T. Philos. Trans. R. Soc. A Math. Phys. Eng. Sci. 2005, 363, 901–912. Shvo, Y.; Czarkie, D.; Rahamim, Y.; Chodosh, D. F. J. Am. Chem. Soc. 1986, 108, 7400–7402. Karvembu, R.; Prabhakaran, R.; Natarajan, K. Coord. Chem. Rev. 2005, 249, 911–918. Quintard, A.; Rodriguez, J. Angew. Chem. Int. Ed. 2014, 53, 4044–4055. Wang, L.; Kanega, R.; Kawanami, H.; Himeda, Y. Chem. Rec. 2017, 17, 1071–1094. Wang, W.-H.; Hull, J. F.; Muckerman, J. T.; Fujita, E.; Himeda, Y. Energy Environ. Sci. 2012, 5, 7923–7926. Burks, D. B.; Vasiliu, M.; Dixon, D. A.; Papish, E. T. J. Phys. Chem. A 2018, 122, 2221–2231. Zhang, T.; Wang, C.; Liu, S.; Wang, J.-L.; Lin, W. J. Am. Chem. Soc. 2014, 136, 273–281. DePasquale, J.; Nieto, I.; Reuther, L. E.; Herbst-Gervasoni, C. J.; Paul, J. J.; Mochalin, V.; Zeller, M.; Thomas, C. M.; Addison, A. W.; Papish, E. T. Inorg. Chem. 2013, 52, 9175–9183. Manbeck, G. F.; Muckerman, J. T.; Szalda, D. J.; Himeda, Y.; Fujita, E. J. Phys. Chem. B. 2015, 119, 7457–7466. Nieto, I.; Livings, M. S.; Sacci, J. B.; Reuther, L. E.; Zeller, M.; Papish, E. T. Organometallics 2011, 30, 6339–6342. Duan, J.; Senger, M.; Esselborn, J.; Engelbrecht, V.; Wittkamp, F.; Apfel, U.-P.; Hofmann, E.; Stripp, S. T.; Happe, T.; Winkler, M. Nat. Commun. 2018, 9, 4726. Sommer, C.; Adamska-Venkatesh, A.; Pawlak, K.; Birrell, J. A.; Rüdiger, O.; Reijerse, E. J.; Lubitz, W. J. Am. Chem. Soc. 2017, 139, 1440–1443. Casey, C. P.; Guan, H. Organometallics 2012, 31, 2631–2638. Rossini, E.; Bochevarov, A. D.; Knapp, E. W. ACS Omega 2018, 3, 1653–1662. Dub, P. A.; Gordon, J. C. ACS Catal. 2017, 7, 6635–6655. Dub, P. A.; Henson, N. J.; Martin, R. L.; Gordon, J. C. J. Am. Chem. Soc. 2014, 136, 3505–3521. Govindarajan, N.; Beks, H.; Meijer, E. J. ACS Catal. 2020, 10, 14775–14781. Siek, S.; Burks, D. B.; Gerlach, D. L.; Liang, G.; Tesh, J. M.; Thompson, C. R.; Qu, F.; Shankwitz, J. E.; Vasquez, R. M.; Chambers, N. S.; Szulczewski, G. J.; Grotjahn, D. B.; Webster, C. E.; Papish, E. T. Organometallics 2017, 36, 1091–1106. Borovik, A. S. Acc. Chem. Res. 2005, 38, 54–61. Yao, W.; Das, S.; DeLucia, N. A.; Qu, F.; Boudreaux, C. M.; Vannucci, A. K.; Papish, E. T. Organometallics 2020, 39, 662–669. Marelius, D. C.; Bhagan, S.; Charboneau, D. J.; Schroeder, K. M.; Kamdar, J. M.; McGettigan, A. R.; Freeman, B. J.; Moore, C. E.; Rheingold, A. L.; Cooksy, A. L.; Smith, D. K.; Paul, J. J.; Papish, E. T.; Grotjahn, D. B. Eur. J. Inorg. Chem. 2014, 2014, 676–689. Das, S.; Nugegoda, D.; Qu, F.; Boudreaux, C. M.; Burrow, P. E.; Figgins, M. T.; Lamb, R. W.; Webster, C. E.; Delcamp, J. H.; Papish, E. T. Eur. J. Inorg. Chem. 2020, 2020, 2709–2717. Dub, P. A.; Gordon, J. C. Nat. Rev. Chem. 2018, 2, 396–408. Pignataro, L.; Gennari, C. Eur. J. Org. Chem. 2020, 2020, 3192–3205. Dogutan, D. K.; McGuire, R., Jr.; Nocera, D. G. J. Am. Chem. Soc. 2011, 133, 9178–9180.

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

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. 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.

469

Lee, C. H.; Dogutan, D. K.; Nocera, D. G. J. Am. Chem. Soc. 2011, 133, 8775–8777. Hashiguchi, B. G.; Bischof, S. M.; Konnick, M. M.; Periana, R. A. Acc. Chem. Res. 2012, 45, 885–898. Dixon, N. A.; McQuarters, A. B.; Kraus, J. S.; Soffer, J. B.; Lehnert, N.; Schweitzer-Stenner, R.; Papish, E. T. Chem. Commun. 2013, 5571–5573. Papish, E. T.; Dixon, N. A.; Kumar, M. Struct. Bonding 2014, 160, 115–150. Agarwal, J.; Shaw, T. W.; Schaefer, I.; Henry, F.; Bocarsly, A. B. Inorg. Chem. 2015, 54, 5285–5294. Labrum, N. S.; Curtin, G. M.; Jakubikova, E.; Caulton, K. G. Chem. Eur. J. 2020, 26, 9547–9555. Heyer, A. J.; Shivokevich, P. J.; Hooe, S. L.; Welch, K. D.; Harman, W. D.; Machan, C. W. Dalton Trans. 2018, 47, 6323–6332. Kumar, M.; Papish, E. T.; Zeller, M.; Hunter, A. D. Dalton Trans. 2011, 40, 7517. Kumar, M.; Papish, E. T.; Zeller, M.; Hunter, A. D. Dalton Trans. 2010, 39, 59–61. Procter, R. J.; Uzelac, M.; Cid, J.; Rushworth, P. J.; Ingleson, M. J. ACS Catal. 2019, 9, 5760–5771. Siek, S.; Dixon, N. A.; Kumar, M.; Kraus, J. S.; Wells, K. R.; Rowe, B. W.; Kelley, S. P.; Zeller, M.; Yap, G. P. A.; Papish, E. T. Eur. J. Inorg. Chem. 2016, 2495–2507. Park, Y. J.; Sickerman, N. S.; Ziller, J. W.; Borovik, A. S. Chem. Commun. 2010, 46, 2584–2586. Hartle, M. D.; Delgado, M.; Gilbertson, J. D.; Pluth, M. D. Chem. Commun. 2016, 52, 7680–7682. Noyori, R.; Ohkuma, T. Angew. Chem. Int. Ed. 2001, 40, 40–73. Ikariya, T.; Murata, K.; Noyori, R. Org. Biomol. Chem. 2006, 4, 393–406. Li, X.; Wu, X.; Chen, W.; Hancock, F. E.; King, F.; Xiao, J. Org. Lett. 2004, 6, 3321–3324. Noyori, R.; Yamakawa, M.; Hashiguchi, S. J. Org. Chem. 2001, 66, 7931–7944. Takehara, J.; Hashiguchi, S.; Fujii, A.; Inoue, S.-i.; Ikariya, T.; Noyori, R. Chem. Commun. 1996, 233–234. Fujii, A.; Hashiguchi, S.; Uematsu, N.; Ikariya, T.; Noyori, R. J. Am. Chem. Soc. 1996, 118, 2521–2522. Hashiguchi, S.; Fujii, A.; Takehara, J.; Ikariya, T.; Noyori, R. J. Am. Chem. Soc. 1995, 117, 7562–7563. Matsumura, K.; Arai, N.; Hori, K.; Saito, T.; Sayo, N.; Ohkuma, T. J. Am. Chem. Soc. 2011, 133, 10696–10699. Terrade, F. G.; Lutz, M.; Van Der Vlugt, J. I.; Reek, J. N. H. Eur. J. Inorg. Chem. 2014, 2014, 1826–1835. John, J. M.; Takebayashi, S.; Dabral, N.; Miskolzie, M.; Bergens, S. H. J. Am. Chem. Soc. 2013, 135, 8578–8584. John, J. M.; Bergens, S. H. Angew. Chem. Int. Ed. 2011, 50, 10377–10380. Hartmann, R.; Chen, P. Angew. Chem. Int. Ed. 2001, 40, 3581–3585. Xie, J.-B.; Xie, J.-H.; Liu, X.-Y.; Kong, W.-L.; Li, S.; Zhou, Q.-L. J. Am. Chem. Soc. 2010, 132, 4538–4539. Xie, J.-B.; Xie, J.-H.; Liu, X.-Y.; Zhang, Q.-Q.; Zhou, Q.-L. Chem. Asian J. 2011, 6, 899–908. Xie, J.-H.; Liu, X.-Y.; Xie, J.-B.; Wang, L.-X.; Zhou, Q.-L. Angew. Chem. Int. Ed. 2011, 50, 7329–7332. Wu, W.; Liu, S.; Duan, M.; Tan, X.; Chen, C.; Xie, Y.; Lan, Y.; Dong, X.-Q.; Zhang, X. Org. Lett. 2016, 18, 2938–2941. Dub, P. A.; Scott, B. L.; Gordon, J. C. Organometallics 2015, 34, 4464–4479. Samec, J. S. M.; Backvall, J.-E.; Andersson, P. G.; Brandt, P. Chem. Soc. Rev. 2006, 35, 237–248. Conley, B. L.; Pennington-Boggio, M. K.; Boz, E.; Williams, T. J. Chem. Rev. 2010, 110, 2294–2312. Weber, M. A.; Ford, P. C. J. Mol. Catal. A: Chem. 2016, 416, 81–87. Lee, H.; Kang, B.; Lee, S.-I.; Hong, S. Synlett 2015, 26, 1077–1080. Afanasenko, A.; Hannah, R.; Yan, T.; Elangovan, S.; Barta, K. ChemSusChem 2019, 12, 3801–3807. Hamid, M. H. S. A.; Slatford, P. A.; Williams, J. M. J. Adv. Synth. Catal. 2007, 349, 1555–1575. Zhang, C.; Zhao, J.-P.; Hu, B.; Shi, J.; Chen, D. Organometallics 2019, 38, 654–664. Pan, H.-J.; Ng, T. W.; Zhao, Y. Chem. Commun. 2015, 51, 11907–11910. Del Grosso, A.; Clarkson, G. J.; Wills, M. Inorg. Chim. Acta 2019, 496, 119043. Dou, X.; Hayashi, T. Adv. Synth. Catal. 2016, 358, 1054–1058. Hurem, D.; Dudding, T. RSC Adv. 2015, 5, 101732–101739. van Slagmaat, C. A. M. R.; De Wildeman, S. M. A. Eur. J. Inorg. Chem. 2017, 2018, 694–702. van Slagmaat, C. A. M. R.; Delgove, M. A. F.; Stouten, J.; Morick, L.; van der Meer, Y.; Bernaerts, K. V.; De Wildeman, S. M. A. Green Chem. 2020, 22, 2443–2458. Li, H.; Qiu, C.; Cao, X.; Lu, Y.; Li, G.; He, X.; Lu, Q.; Chen, K.; Ouyang, P.; Tan, W. ACS Appl. Mater. Interfaces 2019, 11, 15718–15726. Van Buijtenen, J.; Van As, B. A. C.; Meuldijk, J.; Palmans, A. R. A.; Vekemans, J. A. J. M.; Hulshof, L. A.; Meijer, E. W. Chem. Commun. 2006, 105, 3163–3169. Yang, Y.; Guo, J.; Ng, H.; Chen, Z.; Teo, P. Chem. Commun. 2014, 50, 2604–2608. Takahashi, K.; Yamashita, M.; Ichihara, T.; Nakano, K.; Nozaki, K. Angew. Chem. Int. Ed. 2010, 49, 4488–4490. Van As, B. A. C.; Van Buijtenen, J.; Mes, T.; Palmans, A. R. A.; Meijer, E. W. Chem. Eur. J. 2007, 13, 8325–8332. Takahashi, K.; Yamashita, M.; Nozaki, K. J. Am. Chem. Soc. 2012, 134, 18746–18757. Yuki, Y.; Takahashi, K.; Tanaka, Y.; Nozaki, K. J. Am. Chem. Soc. 2013, 135, 17393–17400. Furst, M. R. L.; Korkmaz, V.; Gaide, T.; Seidensticker, T.; Behr, A.; Vorholt, A. J. ChemCatChem 2017, 9, 4319–4323. Cesari, C.; Sambri, L.; Zacchini, S.; Zanotti, V.; Mazzoni, R. Organometallics 2014, 33, 2814–2819. Casey, C. P.; Strotman, N. A.; Beetner, S. E.; Johnson, J. B.; Priebe, D. C.; Guzei, I. A. Organometallics 2006, 25, 1236–1244. Casey, C. P.; Strotman, N. A.; Beetner, S. E.; Johnson, J. B.; Priebe, D. C.; Vos, T. E.; Khodavandi, B.; Guzei, I. A. Organometallics 2006, 25, 1230–1235. Spector, I. C.; Olson, C. M.; Massari, A. M. J. Phys. Chem. C 2016, 120, 24877–24884. Blum, Y.; Czarkie, D.; Rahamim, Y.; Shvo, Y. Organometallics 1985, 4, 1459–1461. Blum, Y.; Shvo, Y. Isr. J. Chem. 1984, 24, 144–148. Casey, C. P.; Singer, S. W.; Powell, D. R.; Hayashi, R. K.; Kavana, M. J. Am. Chem. Soc. 2001, 123, 1090–1100. Casey, C. P.; Johnson, J. B. Can. J. Chem. 2005, 83, 1339–1346. Casey, C. P.; Vos, T. E.; Singer, S. W.; Guzei, I. A. Organometallics 2002, 21, 5038–5046. Casey, C. P.; Johnson, J. B.; Singer, S. W.; Cui, Q. J. Am. Chem. Soc. 2005, 127, 3100–3109. Gusev, D. G.; Spasyuk, D. M. ACS Catal. 2018, 8, 6851–6861. Zhang, X.; Lu, Z.; Foellmer, L. K.; Williams, T. J. Organometallics 2015, 34, 3732–3738. Knolker, H.-J.; Baum, E.; Goesmann, H.; Klauss, R. Angew. Chem. Int. Ed. 1999, 38, 2064–2066. Casey, C. P.; Guan, H. J. Am. Chem. Soc. 2007, 129, 5816–5817. Thorson, M. K.; Klinkel, K. L.; Wang, J.; Williams, T. J. Eur. J. Inorg. Chem. 2009, 2009, 295–302. Pan, H.-J.; Ng, T. W.; Zhao, Y. Org. Biomol. Chem. 2016, 14, 5490–5493. Yan, T.; Feringa, B. L.; Barta, K. Nat. Commun. 2014, 5, 5602–5607. Roudier, T. C. A. Q.a.J. R. M. ACS Catal. 2016, 6, 1–9. Blank, J. H.; Raju, R. K.; Yan, T.; Brothers, E. N.; Darensbourg, M. Y.; Bengali, A. A. Dalton Trans. 2016, 45, 12292–12296. Vailati Facchini, S.; Neudörfl, J.-M.; Pignataro, L.; Cettolin, M.; Gennari, C.; Berkessel, A.; Piarulli, U. ChemCatChem 2017, 9, 1461–1468. Lu, X.; Zhang, Y.; Turner, N.; Zhang, M.; Li, T. Org. Biomol. Chem. 2014, 12, 4361–4371. Ge, H.; Chen, X.; Yang, X. Chem. Commun. 2016, 52, 12422–12425.

470

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. 191. 192. 193. 194. 195. 196. 197. 198. 199. 200. 201. 202. 203. 204. 205. 206. 207. 208. 209. 210. 211. 212.

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

Plank, T. N.; Drake, J. L.; Kim, D. K.; Funk, T. W. Adv. Synth. Catal. 2012, 354, 597–601. Chakraborty, A.; Kinney, R. G.; Krause, J. A.; Guan, H. ACS Catal. 2016, 6, 1–10. von der Höh, A.; Berkessel, A. ChemCatChem 2011, 3, 861–867. Casey, C. P.; Guan, H. J. Am. Chem. Soc. 2009, 131, 2499–2507. Kamitani, M.; Nishiguchi, Y.; Tada, R.; Itazaki, M.; Nakazawa, H. Organometallics 2014, 33, 1532–1535. Berkessel, A.; Reichau, S.; von der Höh, A.; Leconte, N.; Neudörfl, J.r.-M. Organometallics 2011, 30, 3880–3887. Hodgkinson, R.; Del Grosso, A.; Clarkson, G.; Wills, M. Dalton Trans. 2016, 45, 3992–4005. Fleischer, S.; Zhou, S.; Werkmeister, S.; Junge, K.; Beller, M. Chem. Eur. J. 2013, 19, 4997–5003. Zhou, S.; Fleischer, S.; Junge, K.; Beller, M. Angew. Chem. Int. Ed. 2011, 50, 5120–5124. Hopmann, K. H. Chem. Eur. J. 2015, 21, 10020–10030. El-Sepelgy, O.; Brzozowska, A.; Rueping, M. ChemSusChem 2017, 10, 1664–1668. El-Sepelgy, O.; Alandini, N.; Rueping, M. Angew. Chem. 2016, 128, 13800–13803. Hoffmann, F.; Wagler, J.; Roewer, G. Organometallics 2014, 33, 5622–5625. Wu, W.; Solis-Ibarra, D.; Walker, K. L.; Waymouth, R. M. Organometallics 2018, 37, 3298–3302. Kusumoto, S.; Akiyama, M.; Nozaki, K. J. Am. Chem. Soc. 2013, 135, 18726–18729. Kusumoto, S.; Nozaki, K. Nat. Commun. 2015, 6, 6296. Kusumoto, S.; Tatsuki, T.; Nozaki, K. Angew. Chem. 2015, 127, 8578–8581. Ando, H.; Kusumoto, S.; Wu, W.; Nozaki, K. Organometallics 2017, 36, 2317–2322. Higashi, T.; Kusumoto, S.; Nozaki, K. Angew. Chem. Int. Ed. 2021, 60, 2844–2848. Higashi, T.; Ando, H.; Kusumoto, S.; Nozaki, K. J. Am. Chem. Soc. 2019, 141, 2247–2250. Landwehr, A.; Dudle, B.; Fox, T.; Blacque, O.; Berke, H. Chem. Eur. J. 2012, 18, 5701–5714. Raju, S.; van Slagmaat, C. A. M. R.; Lutz, M.; Kleijn, H.; Jastrzebski, J. T. B. H.; Moret, M. E.; Klein Gebbink, R. J. M. Eur. J. Inorg. Chem. 2017, 2017, 741–751. Shima, S.; Pilak, O.; Vogt, S.; Schick, M.; Stagni, M. S.; Meyer-Klaucke, W.; Warkentin, E.; Thauer, R. K.; Ermler, U. Science 2008, 321, 572–575. Lubitz, W.; Ogata, H.; Ruediger, O.; Reijerse, E. Chem. Rev. 2014, 114, 4081–4148. Yang, X.; Hall, M. B. J. Am. Chem. Soc. 2008, 130, 14036–14037. Seo, J.; Ali, A. K.; Rose, M. J. Comments Inorg. Chem. 2014, 34, 103–113. Gerlach, D. L.; Nieto, I.; Herbst-Gervasoni, C. J.; Ferrence, G. M.; Zeller, M.; Papish, E. T. Acta Cryst. Sect. E 2015, E71, 1447–1453. Rawson, J. M.; Winpenny, R. E. P. Coord. Chem. Rev. 1995, 139, 313–374. Yamaguchi, R.; Ikeda, C.; Takahashi, Y.; Fujita, K.-i. J. Am. Chem. Soc. 2009, 131, 8410–8412. Fujita, K.-i.; Tanino, N.; Yamaguchi, R. Org. Lett. 2007, 9, 109–111. Royer, A. M.; Rauchfuss, T. B.; Gray, D. L. Organometallics 2010, 29, 6763–6768. Li, H.; Jiang, J.; Lu, G.; Huang, F.; Wang, Z.-X. Organometallics 2011, 30, 3131–3141. Li, H.; Lu, G.; Jiang, J.; Huang, F.; Wang, Z.-X. Organometallics 2011, 30, 2349–2363. Cooksey, J. P.; Saidi, O.; Williams, J. M. J.; Blacker, A. J.; Marsden, S. P. Tetrahedron 2021, 78, 131785. Gallardo-Villagrán, M.; Rivada-Wheelaghan, O.; Rahaman, S. M. W.; Fayzullin, R. R.; Khusnutdinova, J. R. Dalton Trans. 2020, 49, 12756–12766. Jiang, F.; Achard, M.; Roisnel, T.; Dorcet, V.; Bruneau, C. Eur. J. Inorg. Chem. 2015, 2015, 4312–4317. Himeda, Y.; Onozawa-Komatsuzaki, N.; Sugihara, H.; Arakawa, H.; Kasuga, K. Organometallics 2004, 23, 1480–1483. Himeda, Y.; Onozawa-Komatsuzaki, N.; Sugihara, H.; Kasuga, K. Organometallics 2007, 26, 702–712. Himeda, Y.; Miyazawa, S.; Hirose, T. ChemSusChem 2011, 4, 487–493. Sordakis, K.; Tsurusaki, A.; Iguchi, M.; Kawanami, H.; Himeda, Y.; Laurenczy, G. Chem. Eur. J. 2016, 22, 15605–15608. Conifer, C. M.; Taylor, R. A.; Law, D. J.; Sunley, G. J.; White, A. J. P.; Britovsek, G. J. P. Dalton Trans. 2011, 40, 1031–1033. Conifer, C. M.; Law, D. J.; Sunley, G. J.; Haynes, A.; Wells, J. R.; White, A. J. P.; Britovsek, G. J. P. Eur. J. Inorg. Chem. 2011, 2011, 3511–3522. E.T. Papish, I. Nieto, US Patent WO2013033018A2, (2013). Nacianceno, V. S.; Garralda, M. A.; Matxain, J. M.; Freixa, Z. Organometallics 2020, 39, 1238–1248. Liu, P.; Tung, N. T.; Xu, X.; Yang, J.; Li, F. J. Org. Chem. 2021, 86, 2621–2631. Ertem, M. Z.; Himeda, Y.; Fujita, E.; Muckerman, J. T. ACS Catal. 2016, 6, 600–609. Kawahara, R.; Fujita, K.-i.; Yamaguchi, R. J. Am. Chem. Soc. 2012, 134, 3643–3646. Yao, W.; DeRegnaucourt, A. R.; Shrewsbury, E. D.; Loadholt, K. H.; Silprakob, W.; Brewster, T. P.; Papish, E. T. Organometallics 2020, 39, 3656–3662. Brewster, T. P.; Ou, W. C.; Tran, J. C.; Goldberg, K. I.; Hanson, S. K.; Cundari, T. R.; Heinekey, D. M. ACS Catal. 2014, 4, 3034–3038. Wang, W.-H.; Xu, S.; Manaka, Y.; Suna, Y.; Kambayashi, H.; Muckerman, J. T.; Fujita, E.; Himeda, Y. ChemSusChem 2014, 7, 1976–1983. Fujita, K.-i.; Tanaka, Y.; Kobayashi, M.; Yamaguchi, R. J. Am. Chem. Soc. 2014, 136, 4829–4832. Zeng, G.; Sakaki, S.; Fujita, K.-i.; Sano, H.; Yamaguchi, R. ACS Catal. 2014, 4, 1010–1020. Fujita, K.-i.; Kawahara, R.; Aikawa, T.; Yamaguchi, R. Angew. Chem. 2015, 127, 9185–9188. Kawahara, R.; Fujita, K.-i.; Yamaguchi, R. Angew. Chem. Int. Ed. 2012, 51, 12790–12794. Wang, R.; Ma, J.; Li, F. J. Org. Chem. 2015, 80, 10769–10776. Li, F.; Lu, L.; Liu, P. Org. Lett. 2016, 18, 2580–2583. Roy, B. C.; Chakrabarti, K.; Shee, S.; Paul, S.; Kundu, S. Chem. Eur. J. 2016, 22, 18147–18155. Sahoo, A. R.; Jiang, F.; Bruneau, C.; Sharma, G. V. M.; Suresh, S.; Roisnel, T.; Dorcet, V.; Achard, M. Catal. Sci. Technol. 2017, 7, 3492–3498. Sahoo, A. R.; Lalitha, G.; Murugesh, V.; Bruneau, C.; Sharma, G. V. M.; Suresh, S.; Achard, M. J. Org. Chem. 2017, 82, 10727–10731. Shi, J.; Hu, B.; Gong, D.; Shang, S.; Hou, G.; Chen, D. Dalton Trans. 2016, 45, 4828–4834. Shi, J.; Hu, B.; Ren, P.; Shang, S.; Yang, X.; Chen, D. Organometallics 2018, 37, 2795–2806. Shi, J.; Hu, B.; Chen, X.; Shang, S.; Deng, D.; Sun, Y.; Shi, W.; Yang, X.; Chen, D. ACS Omega 2017, 2, 3406–3416. Deng, D.; Hu, B.; Yang, M.; Chen, D. Organometallics 2018, 37, 3353–3359. de Boer, S. Y.; Korstanje, T. J.; La Rooij, S. R.; Kox, R.; Reek, J. N. H.; van der Vlugt, J. I. Organometallics 2017, 36, 1541–1549. Paul, B.; Chakrabarti, K.; Kundu, S. Dalton Trans. 2016, 45, 11162–11171. Hale, L. V. A.; Szymczak, N. K. ACS Catal. 2018, 8, 6446–6461. Moore, C. M.; Szymczak, N. K. Chem. Commun. 2013, 49, 400–402. Moore, C. M.; Bark, B.; Szymczak, N. K. ACS Catal. 2016, 6, 1981–1990. Dahl, E. W.; Louis-Goff, T.; Szymczak, N. K. Chem. Commun. 2017, 53, 2287–2289. Geri, J. B.; Szymczak, N. K. J. Am. Chem. Soc. 2015, 137, 12808–12814. Geri, J. B.; Ciatti, J. L.; Szymczak, N. K. Chem. Commun. 2018, 54, 7790–7793. Hull, J. F.; Himeda, Y.; Wang, W.-H.; Hashiguchi, B.; Periana, R.; Szalda, D. J.; Muckerman, J. T.; Fujita, E. Nat. Chem. 2012, 4, 383–388. Iguchi, M.; Zhong, H.; Himeda, Y.; Kawanami, H. Chem. Eur. J. 2017, 23, 17788–17793.

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

213. 214. 215. 216. 217. 218. 219. 220. 221. 222. 223. 224. 225. 226. 227. 228. 229. 230. 231. 232. 233. 234. 235. 236. 237. 238. 239. 240. 241. 242. 243. 244. 245. 246. 247. 248. 249. 250. 251. 252. 253. 254. 255. 256. 257. 258. 259. 260. 261. 262. 263. 264. 265. 266. 267. 268. 269. 270. 271. 272. 273. 274. 275. 276. 277. 278. 279. 280.

471

Kanega, R.; Ertem, M. Z.; Onishi, N.; Szalda, D. J.; Fujita, E.; Himeda, Y. Organometallics 2020, 39, 1519–1531. Kanega, R.; Onishi, N.; Szalda, D. J.; Ertem, M. Z.; Muckerman, J. T.; Fujita, E.; Himeda, Y. ACS Catal. 2017, 7, 6426–6429. Kanega, R.; Onishi, N.; Wang, L.; Himeda, Y. ACS Catal. 2018, 8, 11296–11301. Onishi, N.; Laurenczy, G.; Beller, M.; Himeda, Y. Coord. Chem. Rev. 2018, 373, 317–332. Suna, Y.; Ertem, M. Z.; Wang, W.-H.; Kambayashi, H.; Manaka, Y.; Muckerman, J. T.; Fujita, E.; Himeda, Y. Organometallics 2014, 33, 6519–6530. Wang, L.; Ertem, M. Z.; Murata, K.; Muckerman, J. T.; Fujita, E.; Himeda, Y. ACS Catal. 2018, 8, 5233–5239. Wang, W.-H.; Tang, H.-P.; Lu, W.-D.; Li, Y.; Bao, M.; Himeda, Y. ChemCatChem 2017, 9, 3191–3196. Wang, W.-H.; Ertem, M. Z.; Xu, S.; Onishi, N.; Manaka, Y.; Suna, Y.; Kambayashi, H.; Muckerman, J. T.; Fujita, E.; Himeda, Y. ACS Catal. 2015, 5, 5496–5504. Onishi, N.; Xu, S.; Manaka, Y.; Suna, Y.; Wang, W.-H.; Muckerman, J. T.; Fujita, E.; Himeda, Y. Inorg. Chem. 2015, 54, 5114–5123. Xu, S.; Onishi, N.; Tsurusaki, A.; Manaka, Y.; Wang, W.-H.; Muckerman, J. T.; Fujita, E.; Himeda, Y. Eur. J. Inorg. Chem. 2015, 2015, 5591–5594. Suna, Y.; Himeda, Y.; Fujita, E.; Muckerman, J. T.; Ertem, M. Z. ChemSusChem 2017, 10, 4535–4543. Wang, L.; Onishi, N.; Murata, K.; Hirose, T.; Muckerman, J. T.; Fujita, E.; Himeda, Y. ChemSusChem 2017, 10, 1071–1075. Huang, M.; Li, Y.; Liu, J.; Lan, X.-B.; Liu, Y.; Zhao, C.; Ke, Z. Green Chem. 2019, 21, 219–224. Fujita, K.-i.; Tamura, R.; Tanaka, Y.; Yoshida, M.; Onoda, M.; Yamaguchi, R. ACS Catal. 2017, 7226–7230. Thoi, V. S.; Chang, C. J. Chem. Commun. 2011, 47, 6578. Thoi, V. S.; Kornienko, N.; Margarit, C. G.; Yang, P.; Chang, C. J. J. Am. Chem. Soc. 2013, 135, 14413–14424. Kawanami, H.; Iguchi, M.; Himeda, Y. Inorg. Chem. 2020, 59, 4191–4199. Wang, W.-H.; Muckerman, J. T.; Fujita, E.; Himeda, Y. ACS Catal. 2013, 3, 856–860. Fujita, K.-i.; Kawahara, R.; Aikawa, T.; Yamaguchi, R. Angew. Chem. Int. Ed. 2015, 54, 9057–9060. Badiei, Y. M.; Wang, W.-H.; Hull, J. F.; Szalda, D. J.; Muckerman, J. T.; Himeda, Y.; Fujita, E. Inorg. Chem. 2013, 52, 12576–12586. Gerlach, D. L.; Bhagan, S.; Cruce, A. A.; Burks, D. B.; Nieto, I.; Truong, H. T.; Kelley, S. P.; Herbst-Gervasoni, C. J.; Jernigan, K. L.; Bowman, M. K.; Pan, S.; Zeller, M.; Papish, E. T. Inorg. Chem. 2014, 53, 12689–12698. Duan, L.; Manbeck, G. F.; Kowalczyk, M.; Szalda, D. J.; Muckerman, J. T.; Himeda, Y.; Fujita, E. Inorg. Chem. 2016, 55, 4582–4594. Taylor, J. O.; Neri, G.; Banerji, L.; Cowan, A. J.; Hartl, F. Inorg. Chem. 2020, 59, 5564–5578. Smieja, J. M.; Kubiak, C. P. Inorg. Chem. 2010, 49, 9283–9289. Costentin, C.; Passard, G.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 2014, 136, 11821–11829. Costentin, C.; Passard, G.; Robert, M.; Savéant, J.-M. Proc. Natl. Acad. Sci. U. S. A. 2014, 111, 14990–14994. Costentin, C.; Robert, M.; Savéant, J.-M.; Tatin, A. Proc. Natl. Acad. Sci. U. S. A. 2015, 112, 6882–6886. Wilting, A.; Stolper, T.; Mata, R. A.; Siewert, I. Inorg. Chem. 2017, 56, 4176–4185. Wilting, A.; Siewert, I. ChemistrySelect 2018, 3, 4593–4597. Du, J.-P.; Wilting, A.; Siewert, I. Chem. A Eur. J. 2019, 25, 5555–5564. Lacy, D. C.; Gupta, R.; Stone, K. L.; Greaves, J.; Ziller, J. W.; Hendrich, M. P.; Borovik, A. S. J. Am. Chem. Soc. 2010, 132, 12188–12190. Shook, R. L.; Borovik, A. S. Inorg. Chem. 2010, 49, 3646–3660. Betley, T. A.; Wu, Q.; Van Voorhis, T.; Nocera, D. G. Inorg. Chem. 2008, 47, 1849–1861. Taguchi, T.; Gupta, R.; Lassalle-Kaiser, B.; Boyce, D. W.; Yachandra, V. K.; Tolman, W. B.; Yano, J.; Hendrich, M. P.; Borovik, A. S. J. Am. Chem. Soc. 2012, 134, 1996–1999. Gupta, R.; Taguchi, T.; Lassalle-Kaiser, B.; Bominaar, E. L.; Yano, J.; Hendrich, M. P.; Borovik, A. S. Proc. Natl. Acad. Sci. U. S. A. 2015, 112, 5319–5324. MacBeth, C. E.; Golombek, A. P.; Young, V. G.; Yang, C.; Kuczera, K.; Hendrich, M. P.; Borovik, A. S. Science 2000, 289, 938–941. Mukherjee, J.; Lucas, R. L.; Zart, M. K.; Powell, D. R.; Day, V. W.; Borovik, A. S. Inorg. Chem. 2008, 47, 5780–5786. Parsell, T. H.; Yang, M.-Y.; Borovik, A. S. J. Am. Chem. Soc. 2009, 131, 2762–2763. Hill, E. A.; Weitz, A. C.; Onderko, E.; Romero-Rivera, A.; Guo, Y.; Swart, M.; Bominaar, E. L.; Green, M. T.; Hendrich, M. P.; Lacy, D. C.; Borovik, A. S. J. Am. Chem. Soc. 2016, 138, 13143–13146. Barman, S. K.; Jones, J. R.; Sun, C.; Hill, E. A.; Ziller, J. W.; Borovik, A. S. J. Am. Chem. Soc. 2019, 141, 11142–11150. Shook, R. L.; Gunderson, W. A.; Greaves, J.; Ziller, J. W.; Hendrich, M. P.; Borovik, A. S. J. Am. Chem. Soc. 2008, 130, 8888–8889. Sahu, S.; Widger, L. R.; Quesne, M. G.; de Visser, S. P.; Matsumura, H.; Moënne-Loccoz, P.; Siegler, M. A.; Goldberg, D. P. J. Am. Chem. Soc. 2013, 135, 10590–10593. Widger, L. R.; Jiang, Y.; McQuilken, A. C.; Yang, T.; Siegler, M. A.; Matsumura, H.; Moënne-Loccoz, P.; Kumar, D.; de Visser, S. P.; Goldberg, D. P. Dalton Trans. 2014, 43, 7522–7532. Lucas, R. L.; Zart, M. K.; Murkerjee, J.; Sorrell, T. N.; Powell, D. R.; Borovik, A. S. J. Am. Chem. Soc. 2006, 128, 15476–15489. Zinn, P. J.; Sorrell, T. N.; Powell, D. R.; Day, V. W.; Borovik, A. S. Inorg. Chem. 2007, 46, 10120–10132. Cook, S. A.; Ziller, J. W.; Borovik, A. S. Inorg. Chem. 2014, 53, 11029–11035. Park, Y. J.; Cook, S. A.; Sickerman, N. S.; Sano, Y.; Ziller, J. W.; Borovik, A. S. Chem. Sci. 2013, 4, 717–726. Park, Y. J.; Ziller, J. W.; Borovik, A. S. J. Am. Chem. Soc. 2011, 133, 9258–9261. Sickerman, N. S.; Peterson, S. M.; Ziller, J. W.; Borovik, A. S. Chem. Commun. 2014, 50, 2515–2517. Sano, Y.; Lau, N.; Weitz, A. C.; Ziller, J. W.; Hendrich, M. P.; Borovik, A. S. Inorg. Chem. 2017, 56, 14118–14128. Biswas, S.; Lau, N.; Borovik, A. S.; Hendrich, M. P.; Bominaar, E. L. Inorg. Chem. 2019, 58, 9150–9160. Cook, S. A.; Bogart, J. A.; Levi, N.; Weitz, A. C.; Moore, C.; Rheingold, A. L.; Ziller, J. W.; Hendrich, M. P.; Borovik, A. S. Chem. Sci. 2018, 9, 6540–6547. Moore, C. M.; Szymczak, N. K. Dalton Trans. 2012, 41, 7886–7889. Lacy, D. C.; Park, Y. J.; Ziller, J. W.; Yano, J.; Borovik, A. S. J. Am. Chem. Soc. 2012, 134, 17526–17535. Oswald, V. F.; Lee, J. L.; Biswas, S.; Weitz, A. C.; Mittra, K.; Fan, R.; Li, J.; Zhao, J.; Hu, M. Y.; Alp, E. E.; Bominaar, E. L.; Guo, Y.; Green, M. T.; Hendrich, M. P.; Borovik, A. S. J. Am. Chem. Soc. 2020, 142, 11804–11817. Oswald, V. F.; Weitz, A. C.; Biswas, S.; Ziller, J. W.; Hendrich, M. P.; Borovik, A. S. Inorg. Chem. 2018, 57, 13341–13350. Jones, J. R.; Ziller, J. W.; Borovik, A. S. Inorg. Chem. 2017, 56, 1112–1120. Lau, N.; Sano, Y.; Ziller, J. W.; Borovik, A. S. Polyhedron 2017, 125, 179–185. Powell-Jia, D.; Ziller, J. W.; DiPasquale, A. G.; Rheingold, A. L.; Borovik, A. S. Dalton Trans. 2009, 2986. Delgado, M.; Gilbertson, J. D. Chem. Commun. 2017, 53, 11249–11252. Kwon, Y. M.; Delgado, M.; Zakharov, L. N.; Seda, T.; Gilbertson, J. D. Chem. Commun. 2016, 52, 11016–11019. Kendall, A. J.; Zakharov, L. N.; Gilbertson, J. D. Inorg. Chem. 2010, 49, 8656–8658. Cheung, P. M.; Burns, K. T.; Kwon, Y. M.; Deshaye, M. Y.; Aguayo, K. J.; Oswald, V. F.; Seda, T.; Zakharov, L. N.; Kowalczyk, T.; Gilbertson, J. D. J. Am. Chem. Soc. 2018, 140, 17040–17050. Rivas, J. C. M.; Hinchley, S. L.; Metteau, L.; Parsons, S. Dalton Trans. 2006, 2316. Natale, D.; Mareque-Rivas, J. C. Chem. Commun. 2008, 425–437. Garner, D. K.; Fitch, S. B.; McAlexander, L. H.; Bezold, L. M.; Arif, A. M.; Berreau, L. M. J. Am. Chem. Soc. 2002, 124, 9970–9971. Yamaguchi, S.; Wada, A.; Funahashi, Y.; Nagatomo, S.; Kitagawa, T.; Jitsukawa, K.; Masuda, H. Eur. J. Inorg. Chem. 2003, 2003, 4378–4386. Yadav, V.; Siegler, M. A.; Goldberg, D. P. J. Am. Chem. Soc. 2021, 143, 46–52.

472

281. 282. 283. 284. 285. 286. 287. 288. 289. 290. 291. 292. 293. 294. 295. 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. 333. 334. 335. 336. 337. 338. 339. 340. 341. 342. 343. 344. 345. 346. 347. 348. 349.

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

Yadav, V.; Rodriguez, R. J.; Siegler, M. A.; Goldberg, D. P. J. Am. Chem. Soc. 2020, 142, 7259–7264. Wijeratne, G. B.; Bhadra, M.; Siegler, M. A.; Karlin, K. D. J. Am. Chem. Soc. 2019, 141, 17962–17967. Rudzka, K.; Arif, A. M.; Berreau, L. M. J. Am. Chem. Soc. 2006, 128, 17018–17023. Harata, M.; Jitsukawa, K.; Masuda, H.; Einaga, H. J. Am. Chem. Soc. 1994, 116, 10817–10818. Wada, A.; Harata, M.; Hasegawa, K.; Jitsukawa, K.; Masuda, H.; Mukai, M.; Kitagawa, T.; Einaga, H. Angew. Chem. Int. Ed. 1998, 37, 798. Goldcamp, M. J.; Robison, S. E.; Krause Bauer, J. A.; Baldwin, M. J. Inorg. Chem. 2002, 41, 2307–2309. Edison, S. E.; Conklin, S. D.; Kaval, N.; Cheruzel, L. E.; Krause, J. A.; Seliskar, C. J.; Heineman, W. R.; Buchanan, R. M.; Baldwin, M. J. Inorg. Chim. Acta 2008, 361, 947–955. Letko, C. S.; Rauchfuss, T. B.; Zhou, X.; Gray, D. L. Inorg. Chem. 2012, 51, 4511–4520. Thorseth, M. A.; Letko, C. S.; Tse, E. C. M.; Rauchfuss, T. B.; Gewirth, A. A. Inorg. Chem. 2013, 52, 628–634. Moore, C. M.; Quist, D. A.; Kampf, J. W.; Szymczak, N. K. Inorg. Chem. 2014, 53, 3278–3280. Moore, C. M.; Szymczak, N. K. Chem. Commun. 2014, 1–3. Serpas, L.; Baum, R. R.; McGhee, A.; Nieto, I.; Jernigan, K. L.; Zeller, M.; Ferrence, G. M.; Tierney, D. L.; Papish, E. T. Polyhedron 2016, 114, 62–71. Moore, C. M.; Szymczak, N. K. Chem. Sci. 2015, 6, 3373–3377. Dahl, E. W.; Dong, H. T.; Szymczak, N. K. Chem. Commun. 2018, 54, 892–895. Wilson, J. R.; Zeller, M.; Szymczak, N. K. Chem. Commun. 2021, 57, 753–756. Dahl, E. W.; Kiernicki, J. J.; Zeller, M.; Szymczak, N. K. J. Am. Chem. Soc. 2018, 140, 10075–10079. Ford, C. L.; Park, Y. J.; Matson, E. M.; Gordon, Z.; Fout, A. R. Science 2016, 354, 741–743. Matson, E. M.; Park, Y. J.; Fout, A. R. J. Am. Chem. Soc. 2014, 136, 17398–17401. Matson, E. M.; Bertke, J. A.; Fout, A. R. Inorg. Chem. 2014, 53, 4450–4458. Matson, E. M.; Park, Y. J.; Bertke, J. A.; Fout, A. R. Dalton Trans. 2015, 44, 10377–10384. Gordon, Z.; Drummond, M. J.; Bogart, J. A.; Schelter, E. J.; Lord, R. L.; Fout, A. R. Inorg. Chem. 2017, 56, 4852–4863. Gordon, Z.; Miller, T. J.; Leahy, C. A.; Matson, E. M.; Burgess, M.; Drummond, M. J.; Popescu, C. V.; Smith, C. M.; Lord, R. L.; Rodríguez-López, J.; Fout, A. R. Inorg. Chem. 2019, 58, 15801–15811. Drummond, M. J.; Miller, T. J.; Ford, C. L.; Fout, A. R. ACS Catal. 2020, 10, 3175–3182. Drummond, M. J.; Ford, C. L.; Gray, D. L.; Popescu, C. V.; Fout, A. R. J. Am. Chem. Soc. 2019, 141, 6639–6650. Park, Y. J.; Matson, E. M.; Nilges, M. J.; Fout, A. R. Chem. Commun. 2015, 51, 5310–5313. Matson, E. M.; Gordon, Z.; Lin, B.; Nilges, M. J.; Fout, A. R. Dalton Trans. 2014, 43, 16992–16995. Yamagishi, H.; Nabeya, S.; Ikariya, T.; Kuwata, S. Inorg. Chem. 2015, 54, 11584–11586. Trofimenko, S. du Pont de Nemours; E. I., and Co., 1966; pp 1–10 Trofimenko, S. J. Am. Chem. Soc. 1967, 89, 3170–3177. Janiak, C. Chem. Ber. 1994, 127, 1379–1385. Caballero, A.; Despagnet-Ayoub, E.; Díaz-Requejo, M. M.; Díaz-Rodríguez, A.; González-Núñez, M. E.; Mello, R.; Muñoz, B. K.; Ojo, W.-S.; Asensio, G.; Etienne, M.; Pérez, P. J. Science 2011, 332, 835–838. Conde, A.; Vilella, L.; Balcells, D.; Díaz-Requejo, M. M.; Lledós, A.; Pérez, P. J. J. Am. Chem. Soc. 2013, 135, 3887–3896. Kou, X.; von Dias, H. V. R. Dalton Trans. 2009, 7529. Kou, X.; Wu, J.; Cundari, T. R.; Dias, H. V. R. Dalton Trans. 2009, 915–917. Oseback, S. N.; Shim, S. W.; Kumar, M.; Greer, S. M.; Gardner, S. R.; Lemar, K. M.; DeGregory, P. R.; Papish, E. T.; Tierney, D. L.; Zeller, M.; Yap, G. P. A. Dalton Trans. 2012, 41, 2774. Papish, E. T.; Donahue, T. M.; Wells, K. R.; Yap, G. P. A. Dalton Trans. 2008, 2923. Gardner, S. R.; Papish, E. T.; Monillas, W. H.; Yap, G. P. A. J. Inorg. Biochem. 2008, 102, 2179–2183. Bongiovanni, J. L.; Rowe, B. W.; Fadden, P. T.; Taylor, M. T.; Wells, K. R.; Kumar, M.; Papish, E. T.; Yap, G. P. A.; Zeller, M. Inorg. Chim. Acta 2010, 363, 2163–2170. Siek, S.; Dixon, N. A.; Papish, E. T. Inorg. Chim. Acta 2017, 459, 80–86. Kumar, M.; Dixon, N. A.; Merkle, A. C.; Zeller, M.; Lehnert, N.; Papish, E. T. Inorg. Chem. 2012, 51, 7004–7006. Kumar, M.; DePasquale, J.; White, N. J.; Zeller, M.; Papish, E. T. Organometallics 2013, 32, 2135–2144. Nazemi, A.; Cundari, T. R. Organometallics 2019, 38, 3521–3531. Semwal, S.; Choudhury, J. ACS Catal. 2016, 6, 2424–2428. Mironov, O. A.; Bischof, S. M.; Konnick, M. M.; Hashiguchi, B. G.; Ziatdinov, V. R.; Goddard, W. A.; Ahlquist, M.; Periana, R. A. J. Am. Chem. Soc. 2013, 135, 14644–14658. Konnick, M. M.; Bischof, S. M.; Yousufuddin, M.; Hashiguchi, B. G.; Ess, D. H.; Periana, R. A. J. Am. Chem. Soc. 2014, 136, 10085–10094. DuBois, D. L.; Bullock, R. M. Eur. J. Inorg. Chem. 2011, 1017–1027. DuBois, D. L. Inorg. Chem. 2014, 53, 3935–3960. Tronic, T. A.; Rakowski DuBois, M.; Kaminsky, W.; Coggins, M. K.; Liu, T.; Mayer, J. M. Angew. Chem. Int. Ed. 2011, 50, 10936–10939. Klug, C. M.; Dougherty, W. G.; Kassel, W. S.; Wiedner, E. S. Organometallics 2019, 38, 1269–1279. Wiedner, E. S.; Roberts, J. A. S.; Dougherty, W. G.; Kassel, W. S.; DuBois, D. L.; Bullock, R. M. Inorg. Chem. 2013, 52, 9975–9988. Wiedner, E. S.; Appel, A. M.; DuBois, D. L.; Bullock, R. M. Inorg. Chem. 2013, 52, 14391–14403. Burgess, S. A.; Grubel, K.; Appel, A. M.; Wiedner, E. S.; Linehan, J. C. Inorg. Chem. 2017, 56, 8580–8589. Prokopchuk, D. E.; Wiedner, E. S.; Walter, E. D.; Popescu, C. V.; Piro, N. A.; Kassel, W. S.; Bullock, R. M.; Mock, M. T. J. Am. Chem. Soc. 2017, 139, 9291–9301. Rakowski Dubois, M.; Dubois, D. L. Chem. Soc. Rev. 2008, 38, 62. Appel, A. M.; Bercaw, J. E.; Bocarsly, A. B.; Dobbek, H.; DuBois, D. L.; Dupuis, M.; Ferry, J. G.; Fujita, E.; Hille, R.; Kenis, P. J. A.; Kerfeld, C. A.; Morris, R. H.; Peden, C. H. F.; Portis, A. R.; Ragsdale, S. W.; Rauchfuss, T. B.; Reek, J. N. H.; Seefeldt, L. C.; Thauer, R. K.; Waldrop, G. L. Chem. Rev. 2013, 113, 6621–6658. Umehara, K.; Kuwata, S.; Ikariya, T. J. Am. Chem. Soc. 2013, 135, 6754–6757. Cook, B. J.; Chen, C.-H.; Pink, M.; Lord, R. L.; Caulton, K. G. Inorg. Chim. Acta 2016, 451, 82–91. Labrum, N. S.; Caulton, K. G. Dalton Trans. 2019, 48, 11642–11646. Yoshinari, A.; Tazawa, A.; Kuwata, S.; Ikariya, T. Chem. Asian J. 2012, 7, 1417–1425. Toda, T.; Suzuki, S.; Kuwata, S. Chem. Commun. 2019, 55, 1028–1031. Nakahara, Y.; Toda, T.; Matsunami, A.; Kayaki, Y.; Kuwata, S. Chem. Asian J. 2017, 13, 73–80. Toda, T.; Kuwata, S.; Ikariya, T. Z. Anorg. Allg. Chem. 2015, 641, 2135–2139. Cook, B. J.; Pink, M.; Pal, K.; Caulton, K. G. Inorg. Chem. 2018, 57, 6176–6185. Cook, B. J.; Chen, C.-H.; Pink, M.; Caulton, K. G. Dalton Trans. 2018, 47, 2052–2060. Polezhaev, A. V.; Chen, C.-H.; Losovyj, Y.; Caulton, K. G. Chem. Eur. J. 2017, 23, 8039–8050. Cabelof, A. C.; Carta, V.; Chen, C.-H.; Caulton, K. G. Dalton Trans. 2020, 49, 7891–7896. Toda, T.; Yoshinari, A.; Ikariya, T.; Kuwata, S. Chem. Eur. J. 2016, 22, 16675–16683. Toda, T.; Kuwata, S. J. Organomet. Chem. 2020, 917, 121270. Toda, T.; Saitoh, K.; Yoshinari, A.; Ikariya, T.; Kuwata, S. Organometallics 2017, 36, 1188–1195.

Proton Responsive and Hydrogen Bonding Ligands in Organometallic Chemistry

350. 351. 352. 353. 354. 355. 356. 357. 358. 359. 360. 361. 362. 363. 364. 365. 366. 367. 368. 369. 370. 371. 372. 373. 374. 375. 376. 377. 378. 379. 380. 381. 382. 383. 384. 385. 386. 387. 388. 389. 390. 391. 392. 393. 394. 395. 396. 397. 398. 399. 400.

473

Araki, K.; Kuwata, S.; Ikariya, T. Organometallics 2008, 27, 2176–2178. Kuwata, S.; Ikariya, T. Chem. Commun. 2014, 50, 14290–14300. Dutta, I.; Yadav, S.; Sarbajna, A.; De, S.; Hölscher, M.; Leitner, W.; Bera, J. K. J. Am. Chem. Soc. 2018, 140, 8662–8666. Kashiwame, Y.; Kuwata, S.; Ikariya, T. Organometallics 2012, 31, 8444–8455. Kashiwame, Y.; Kuwata, S.; Ikariya, T. Chem. Eur. J. 2010, 16, 766–770. Colby, D. A.; Bergman, R. G.; Ellman, J. A. Chem. Rev. 2010, 110, 624–655. Lewis, J. C.; Bergman, R. G.; Ellman, J. A. Acc. Chem. Res. 2008, 41, 1013–1025. Lewis, J. C.; Berman, A. M.; Bergman, R. G.; Ellman, J. A. J. Am. Chem. Soc. 2008, 130, 2493–2500. Lewis, J. C.; Wiedemann, S. H.; Bergman, R. G.; Ellman, J. A. Org. Lett. 2004, 6, 35–38. Tan, K. L.; Bergman, R. G.; Ellman, J. A. J. Am. Chem. Soc. 2001, 123, 2685–2686. Tan, K. L.; Bergman, R. G.; Ellman, J. A. J. Am. Chem. Soc. 2002, 124, 13964–13965. Tan, K. L.; Bergman, R. G.; Ellman, J. A. J. Am. Chem. Soc. 2002, 124, 3202–3203. Tan, K. L.; Park, S.; Ellman, J. A.; Bergman, R. G. J. Org. Chem. 2004, 69, 7329–7335. Miranda-Soto, V.; Grotjahn, D. B.; Cooksy, A. L.; Golen, J. A.; Moore, C. E.; Rheingold, A. L. Angew. Chem. Int. Ed. 2011, 50, 631–635. Miranda-Soto, V.; Grotjahn, D. B.; DiPasquale, A. G.; Rheingold, A. L. J. Am. Chem. Soc. 2008, 130, 13200–13201. Hahn, F. E.; Naziruddin, A. R.; Hepp, A.; Pape, T. Organometallics 2010, 29, 5283–5288. Eguillor, B.; Esteruelas, M. A.; Garcia-Raboso, J.; Olivan, M.; Onate, E.; Pastor, I. M.; Penafiel, I.; Yus, M. Organometallics 2011, 30, 1658–1667. Balof, S. L.; P’Pool, S. J.; Berger, N. J.; Valente, E. J.; Shiller, A. M.; Schanz, H.-J. Dalton Trans. 2008, 5791–5799. He, F.; Danopoulos, A. A.; Braunstein, P. Organometallics 2016, 35, 198–206. Aznarez, F.; Sanz Miguel, P. J.; Tan, T. T. Y.; Hahn, F. E. Organometallics 2016, 35, 410–419. Gomez-Lopez, J. L.; Chávez, D.; Parra-Hake, M.; Royappa, A. T.; Rheingold, A. L.; Grotjahn, D. B.; Miranda-Soto, V. Organometallics 2016, 35, 3148–3153. Norris, M. R.; Flowers, S. E.; Mathews, A. M.; Cossairt, B. M. Organometallics 2016, 35, 2778–2781. Flowers, S. E.; Johnson, M. C.; Pitre, B. Z.; Cossairt, B. M. Dalton Trans. 2018, 47, 1276–1283. Peris, E. Chem. Rev. 2018, 118, 9988–10031. Kuwata, S.; Hahn, F. E. Chem. Rev. 2018, 118, 9642–9677. Jahnke, M. C.; Hahn, F. E. Chem. Lett. 2015, 44, 226–237. Kuwata, S.; Ikariya, T. Chem. Eur. J. 2011, 17, 3542–3556. Fallahpour, R.-A.; Neuburger, M.; Zehnder, M. New J. Chem. 1999, 23, 53–61. Chambers, J.; Eaves, B.; Parker, D.; Claxton, R.; Ray, P. S.; Slattery, S. J. Inorg. Chim. Acta 2006, 359, 2400–2406. Fallahpour, R.-A.; Neuburger, M.; Zehnder, M. Polyhedron 1999, 18, 2445–2454. Burks, D. B.; Davis, S.; Lamb, R. W.; Liu, X.; Rodrigues, R. R.; Liyanage, N. P.; Sun, Y.; Webster, C. E.; Delcamp, J. H.; Papish, E. T. Chem. Commun. 2018, 54, 3819–3822. Sheng, M.; Jiang, N.; Gustafson, S.; You, B.; Ess, D. H.; Sun, Y. Dalton Trans. 2015, 44, 16247–16250. Das, S.; Rodrigues, R. R.; Lamb, R. W.; Qu, F.; Reinheimer, E.; Boudreaux, C. M.; Webster, C. E.; Delcamp, J. H.; Papish, E. T. Inorg. Chem. 2019, 58, 8012–8020. Boudreaux, C. M.; Liyanage, N. P.; Shirley, H.; Siek, S.; Gerlach, D. L.; Qu, F.; Delcamp, J. H.; Papish, E. T. Chem. Commun. 2017, 53, 11217–11220. Rodrigues, R. R.; Boudreaux, C. M.; Papish, E. T.; Delcamp, J. H. ACS Appl. Energy Mater. 2019, 2, 37–46. Dai, Z.; Luo, Q.; Cong, H.; Zhang, J.; Peng, T. New J. Chem. 2017, 41, 3055–3060. Schwartz, T. M.; Burnett, M. E.; Green, K. N. Dalton Trans. 2020, 49, 2356–2363. Ben-Ari, E.; Leitus, G.; Shimon, L. J. W.; Milstein, D. J. Am. Chem. Soc. 2006, 128, 15390–15391. Zhang, J.; Leitus, G.; Ben-David, Y.; Milstein, D. J. Am. Chem. Soc. 2005, 127, 10840–10841. Iron, M. A.; Ben-Ari, E.; Cohen, R.; Milstein, D. Dalton Trans. 2009, 9433–9439. Li, J.; Shiota, Y.; Yoshizawa, K. J. Am. Chem. Soc. 2009, 131, 13584–13585. Gunanathan, C.; Milstein, D. Acc. Chem. Res. 2011, 44, 588–602. Feller, M.; Diskin-Posner, Y.; Shimon, L. J. W.; Ben-Ari, E.; Milstein, D. Organometallics 2012, 31, 4083–4101. Langer, R.; Fuchs, I.; Vogt, M.; Balaraman, E.; Diskin-Posner, Y.; Shimon, L. J. W.; Ben-David, Y.; Milstein, D. Chem. Eur. J. 2013, 19, 3407–3414. Gellrich, U.; Khusnutdinova, J. R.; Leitus, G. M.; Milstein, D. J. Am. Chem. Soc. 2015, 137, 4851–4859. Hashiguchi, B. G.; Young, K. J. H.; Yousufuddin, M.; Goddard, W. A., III; Periana, R. A. J. Am. Chem. Soc. 2010, 132, 12542–12545. Zuo, W.; Lough, A. J.; Li, Y. F.; Morris, R. H. Science 2013, 342, 1080–1083. Zuo, W.; Tauer, S.; Prokopchuk, D. E.; Morris, R. H. Organometallics 2014, 33, 5791–5801. Mikhailine, A. A.; Maishan, M. I.; Lough, A. J.; Morris, R. H. J. Am. Chem. Soc. 2012, 134, 12266–12280. Kounalis, E.; Lutz, M.; Broere, D. L. J. Chem. Eur. J. 2019, 25, 13280–13284. Shimoda, T.; Morishima, T.; Kodama, K.; Hirose, T.; Polyansky, D. E.; Manbeck, G. F.; Muckerman, J. T.; Fujita, E. Inorg. Chem. 2018, 57, 5486–5498.

1.15

Introduction to the Organometallic Chemistry of Carbon Dioxide

Charles W Machan, Department of Chemistry, University of Virginia, Charlottesville, VA, United States © 2022 Elsevier Ltd. All rights reserved.

1.15.1 Introduction 1.15.2 Properties of CO2 as a molecule 1.15.3 M–CO2 complexes 1.15.3.1 Insertion into M–H 1.15.3.2 Insertion into M–C 1.15.3.3 Insertion into M–N 1.15.3.4 Insertion into M–O 1.15.3.5 Metathesis 1.15.4 Homogeneous catalysis with CO2 1.15.4.1 Hydrogenation 1.15.4.2 Electrochemical reduction 1.15.4.3 Carboxylation 1.15.4.4 CO2 copolymerization 1.15.5 Conclusions and perspective Acknowledgment References

1.15.1

474 474 475 478 480 481 482 483 485 485 487 491 491 493 494 494

Introduction

Atmospheric concentrations of carbon dioxide (CO2) have continually risen over the past 200 years, driven by exponential increases in energy demand and corresponding anthropogenic emissions. In the atmosphere, CO2 is a strong absorber of electromagnetic radiation in regions of the spectrum that do not overlap with other small molecules in the atmosphere, e.g., H2O and N2. This “greenhouse effect” contributes to an increased retention of the light energy that strikes Earth from the sun. This retained energy has had consequences on average global temperatures, as well as on the intensity and severity of extreme weather events, a phenomenon widely known as climate change. Remediation strategies involving CO2 capture and conversion will continue to be of essential significance as long as fossil fuel combustion remains a large part of society’s energy portfolio. The thermodynamic stability of CO2 represents a fundamental challenge for selective activation and conversion. For instance, in the classical Kolbe-Schmitt industrial process, the uncatalyzed, direct, reaction of sodium phenoxide with CO2 requires high temperatures and pressures because of its sizable inherent energy barriers.1 For lab-scale preparations,2 the reaction of organolithium precursors or Grignard reagents with CO2 is used to prepare carboxylic acids. In all cases, the highly delocalized OCO p manifold contributes to the thermodynamic stability of CO2, often requiring more forcing conditions to achieve activation. It has been recently estimated that CO2 is an electrophile comparable to benzaldehyde, but that the lack of sufficient thermodynamic driving force results in limited activity with many nucleophiles.3 The ongoing rise of atmospheric CO2 concentrations and anthropogenic emissions sustains interest in using CO2 as a feedstock4–9 to generate more sustainable fuels and commodity chemicals, either by the direct reduction10–12 or through its use as a co-substrate.13–16 If renewable energy sources (e.g., wind, solar) were used to drive the direct reduction of CO2, commodity chemical- and fuel-producing reactions could circularize CO2 emissions.9,17–20 If successful, this strategy would have the dual benefit of diminishing (or reversing) the current level of negative environmental impact and could supplant the use of non-renewable hydrocarbon sources.10,21 In this article, the fundamental properties of CO2 as a ligand in organometallic chemistry and surveys of stoichiometric and catalytic reactions are described. Readers interested in more detail on specific aspects of this overview or heterogeneous processes are directed to more comprehensive reviews.10,13,22–36

1.15.2

Properties of CO2 as a molecule

In its ground state, CO2 has a linear geometry and is non-polar. The CdO bond lengths are short (1.17 A˚ ) compared to other types of carbon-oxygen double bonds (e.g., benzoquinone C–O ¼ 1.223 A˚ ).37,38 This shortening is the consequence of extended p-symmetric molecular orbitals bridging all three atoms, in contrast to the ground state of allene, in which the two CdC double bonds are perpendicular to one another (Fig. 1).39 It is the significant energetic expense of altering the linear geometry CO2 molecule to bent configurations (diminishing its excellent p-overlap) that contributes to its kinetic inertness.37 This stability requires either a strong nucleophile or a bifunctional strategy, where electrophiles interact with one or both of the O atoms and a nucleophile engages the C atom.40–42

474

Comprehensive Organometallic Chemistry IV

https://doi.org/10.1016/B978-0-12-820206-7.00063-9

Introduction to the Organometallic Chemistry of Carbon Dioxide

475

Fig. 1 Comparison of the p-bonding molecular orbitals for CO2 (A) and allene (B).

The direct one-electron reduction of CO2 to the radical anion CO2 − requires a potential of −1.9 V vs. NHE under aprotic conditions to achieve the necessary geometric distortion.37,43 As a result of this required energy input, most reduction reactions involving the activation and conversion of CO2 require highly reducing conditions. Assessing the orbital changes upon geometric distortion of CO2 from linear to bent emphasizes this point (Fig. 2): the LUMO loses its degeneracy as the “in-plane” orbital drops in energy, making reduction more facile.

1.15.3

M–CO2 complexes

The general structural types of mono- and bimetallic CO2 coordination complexes are summarized in Fig. 3. Bridging complexes are known to occur in both homo- and heterobimetallic complexes. There exist several excellent summaries on structurally characterized species from this series.34,44–48 Crystallographic, NMR, and IR data are the most common methods for determining CO2 coordination modes in molecular species. The accessibility of CO2 containing the carbon isotopologue, 13C, and or the oxygen isotopologue, 18O, allows for a variety of IR and NMR spectroscopic studies, which can rigorously establish coordination mode and stereochemistry. From the consideration of a simple harmonic oscillator approximation, one expects that increasing the mass of the carbon and oxygen atoms should shift IR stretching modes related to metal-coordinated CO2 to lower energy. These shifts will in turn be dependent on the symmetry of the stretching modes, which are dictated by which of the C and O atoms are coordinated to the metal center(s) and the overall symmetry of the resultant complex. In NMR spectroscopy, the direct interrogation of 13C or the monitoring of its magnetic coupling to other NMR active nuclei can provide useful structural information. The nature

Fig. 2 Qualitative molecular orbital diagram comparing the linear (A) and bent (B) forms of CO2.

476

Introduction to the Organometallic Chemistry of Carbon Dioxide

Fig. 3 Common CO2 coordination modes with one and two metal centers.36

and type of CO2 coordination modes is often connected to specific reaction products in mechanistic studies, which will be elaborated in subsequent sections.49–51 The interaction of CO2 with a transition metal center (M) was first reported by Vol0 pin et al. in 1969,52 who described a weak IR stretch at 1630 cm−1 appearing after Rh(PPh3)3Cl was treated with CO2 (Ph ¼ phenyl). According to microanalysis, the authors suggested that the best description was (PPh3)3RhCl•CO2•(PPh3)2RhCl, for which no additional structural information concerning geometry or binding mode was specified. This was followed by a report from Jolly et al. in 1971,53 in which a dimeric Ni species of the general formula [Ni(PCy3)2]2(CO2) (Cy ¼ cyclohexyl) was proposed based on elemental analysis and an observed IR band of 1735 cm−1 (KBr). The first definitive structural characterization of a M-CO2 complex was obtained by Aresta et al. in 1975.54 The structure was unambiguously determined to be [Ni(Z2-CO2)(PCy3)2] by X-ray crystallography (Fig. 4); IR characterization in a Nujol mull showed a sharp band at 1740 cm−1. Structurally identical samples could be obtained by treating [Ni(PCy)3)3] or [Ni(PCy3)2]2(N2) with CO2 in toluene or isolated from a mixture of products when [NiBr2(PCy3)2] was reduced with sodium sand in toluene under a CO2 atmosphere. If the [Ni(Z2-CO2)(PCy3)2] complex was dissolved in light petroleum ether, partial CO2 release was observed and a solid was obtained, which had similar combustion analysis and IR characterization data to the earlier report from Jolly et al.53 Notably, the O–C–O angle of the Ni-bound CO2 of 133 reported by Aresta et al.54 is very close to the angle determined for the radical anion CO2 − of 134 .56–58 This suggests that the bonding interaction between Ni and the CO2 fragment has resulted in a partial reduction of CO2. A more complete structural model was reported by Dohring et al. in 1985 (Fig. 4).55 The initial report by Aresta et al. was followed by a Nb(Z2-CO2) adduct from Bristow et al. in 1981 with comparable metrical parameters for the bent CO2, but with longer distances to the Nb center (Fig. 5A).59 A 13C NMR chemical shift of 200.5 ppm was

Fig. 4 Complete structural model obtained from X-ray crystallographic studies of [Ni(Z2-CO2)(PCy3)2] by Dohring et al. in 1985 (CCDC 1135653),55 following the initial report by Aresta et al. in 1975 showing connectivity of the molecular core (CCDC 1121827).54 H atoms omitted for clarity; key bonding metrics: C1–O1 1.210 (4) A˚ ; C1–O2 1.257(5) A˚ ; Ni1–C1 1.857(4) A˚ ; Ni1–O2 1.966(3) A˚ ; O1–C1–O2 136.2(3) .

Introduction to the Organometallic Chemistry of Carbon Dioxide

477

Fig. 5 (A) Structural model for [Nb(Z5-MeCp)2(CH2SiMe3)(Z2-CO2)], CCDC 1105356. H atoms omitted for clarity; key bonding metrics: Nb–C1 2.144(7) A˚ , Nb–O1 2.173(4) A˚ , C1–O1 1.283(8) A˚ , and C1–O2 1.216(8) A˚ ; O1–C1–O2 132.4(7) . (B) Structural model highlighting part of the repetitive subunit for form B of [Co(n-Pr-salen)K(CO2)(THF)]n metallopolymer, CCDC 1148379. H atoms omitted for clarity; key bonding metrics: Co–C1 2.00(2) A˚ , C1–O3 1.20(2) A˚ ; C1–O4 1.24 (2) A˚ ; K1–O3 2.74(1) A˚ ; K1–O4 2.66(1) A˚ ; K2–O3 2.68(1) A˚ ; O3–C1–O4 132.0(15) . (C) Structural model of Rh(diars)2(Cl)(Z1-CO2), CCDC 1120054. H atoms and CO2 positional disorder omitted for clarity; key bonding metrics: Rh–C1 2.063(15) A˚ , C1–O1 1.201(18) A˚ ; C1–O2 1.252(18) A˚ ; O1–C1–O2 128(2) .

reported for the CO2 adduct recorded in CDCl3 solution, with an IR band at 1695 cm−1 assigned to a C–O stretch. Following this report, a multimetallic CO2 adduct was reported by Gambarotta et al. in 1982 by exposing Co(salen) derivatives reduced with alkali metals (Li, Na, K, Cs) to CO2 in a THF solution.60 One of these crystallized in a repetitive polymeric unit, with close contacts between two unique K coordination environments and a Co center (Fig. 5B). A vibrational mode centered at 1650 cm−1 was reported for the bulk isolated material in a Nujol mull. Notably, structural evidence of the first unsupported Z1-CO2 complex was reported in 1983 by Calabrese et al. from DuPont, who disclosed the isolation of a Rh complex, Rh(diars)2(Cl)(Z1-CO2), where diars ¼ 1,2-bis(dimethylarsino)benzene (Fig. 5C).61 Prior to the 1983 report, DuPont had been issued a patent in 197662 for a series of CO2 complexes made from trialkylphosphino Ni(0), Pd(0), and Pt(0) complexes, Rh(I) and Ir(I) bis(diphosphine) compounds, and a single Ir(I) bis(diarsine) complex, providing only IR and microanalysis in support of the proposed structures. The observation of CO2 complexation with the Ir(I) complexes was reported the following year, although X-ray structural evidence was not obtained.63 The 1983 report by Calabrese et al. went on to describe the formation of CO2 adducts from square planar Rh(I) and Ir(I) complexes with mono- and bidentate phosphines, as well as the structurally characterized bis(diarsine) complex. It is worth noting that elevated CO2 pressures were required for the Rh(I) and Ir(I) derivatives.62,63 The structurally characterized species in Fig. 5C was obtained at CO2 pressures of 0.34 atm.61 Exposing the Ir(I) adduct species, Ir(dmpe)2(Cl)(Z1-CO2) (dmpe ¼ 1,1-bis(dimethylphosphino)ethane), to methyl triflate resulted in the alkylation of one of the O atoms; alternatively, binding of the Lewis acid BPh3 also occurred at one of the O atoms. In the ensuing years, a variety of coordination modes and stable complexes were reported.64–68 It is worth mentioning an unusual group of Mo-based complexes, which coordinated two equivalents of CO2 in trans positions relative to one another. In an initial report by Chatt et al.,69 it was established that exposing cis-Mo(N2)2(PMe2Ph)4 with CO2 produced an unstable compound which was proposed to be of the general structural formula “Mo(CO2)2(PMe2Ph)4,” which served as an intermediate to a dimer structure containing a terminal carbonyl group on each Mo atom and two bridging carbonate ligands. The net reaction, two equiv. of CO2 and two electrons generating one equiv. each of CO and carbonate (CO2− 3 ), is also known as reductive disproportionation. NMR and IR data for an analog of the dimeric decomposition product [Mo(CO3)(CO)(PMe3)]2 was obtained later,64 although the putative bis-CO2 intermediate could not be fully characterized. Later work by Carmona and co-workers used analogous Mo-based dinitrogen precursors to generate and characterize a series of these complexes, verifying that the CO2 ligands were bound trans to one another in an Z2 fashion (Fig. 6).64,68,70–72 Interestingly, the crystallographically characterized examples showed that the two Z2-CO2 ligands are orthogonal to one another, suggesting the participation of separate p-symmetric d orbitals in the two bonding interactions. Note that in the bent form (Fig. 2), CO2 presents p-symmetric anti-bonding orbitals which can be populated by donation from the Mo metal center. The first structural evidence of the end-on coordination mode of CO2 was obtained with a uranium complex reported by Meyer and co-workers.73 Previously, this coordination mode had been proposed only based on the results of spectroscopic studies. The authors reported that exposure of a toluene solution with a six-coordinate U(III) complex containing an electron-rich tris(aryloxide) motif to CO2 resulted in rapid discoloration. Crystallographic studies showed that a CO2 adduct had formed, with a UdO bond distance of 2.351(3) A˚ and an O–C–O angle of 178.0(3) ; the proximal carbon-oxygen bond is impacted by U coordination, contracting to 1.122(4) A˚ relative to the distal bond distance of 1.277(4) A˚ . The most intense vibrational absorption corresponding to 12CO2 appears at 2128 cm−1 for the crystalline solid in Nujol mull. In total, experimental evidence suggested formal oxidation of U(III) to U(IV), with a concomitant one-electron reduction of CO2. Despite the formal reduction of coordinated CO2, it is linear rather than bent, which the authors proposed was enforced by the sterically encumbering adamantyl substituents on the ligand framework.

478

Introduction to the Organometallic Chemistry of Carbon Dioxide

Fig. 6 Structural model of trans,mer-[Mo(CO2)2(PMe3)3(CNiPr)] (iPr ¼ isopyropyl). Important bond distances and angles: Mo–C(CO2) 2.105(10) A˚ (av.), Mo–O 2.147(7) A˚ (average), Mo–C1 2.07(2) A˚ , Mo–P1 2.489(4) A˚ , Mo–P2 2.535(4) A˚ , Mo–P3 2.548(5) A˚ ; O–C–O 133.5(10) (average). CCDC 1128177.68

Fig. 7 Structural model of CO2 in an end-on coordination mode in a molecular V complex, V(ONNO)(OH)(Z1-OCO), where ONNO2− ¼ [2,4-Me2-2-(O)C6H2CH2]2N(CH2)2NMe2. Ellipsoids at 50%, selected distances and angles: V1–O1 2.056(3) A˚ , V1–O3 1.915(2) A˚ , V1–O4]1.920(2) A˚ , V1–O5 2.163(3) A˚ , V1–N1 2.143(3) A˚ , V1–N2 2.263(3) A˚ , V1–O1–C1 159.3(3) , O1–V1–N2 174.35(11) , O5–V1–N1 174.57(12) and O4–V1–O3 174.59(10) . CCDC 1862098.74

More recently, a transition metal-based example of end-on coordination of CO2 through the O atom was reported by Viasus et al. (Fig. 7).74 Exposing a V(II) complex V(ONNO)(TMEDA), where TMEDA ¼ tetramethylethylenediamine and ONNO2− ¼ [2,4-Me2-2-(O)-C6H2CH2]2N(CH2)2NMe2), to 1 atm of CO2 at −10  C resulted in the formation of a formally V(III) complex with CO2 in an end-on coordination mode and a terminal hydroxy ligand, V(ONNO)(OH)(Z1-OCO), with concomitant loss of TMEDA. CO was observed in the reaction headspace and isotopic labeling studies confirmed that CO and the hydroxy O atom both arose from a second equivalent of CO2, which underwent reduction under the reaction conditions; toluene was proposed to be the source of H from isotopic labeling experiments. Interestingly, the end-on CO2 complex was stable to a temperature of 70  C in toluene, at which point CO2 dissociation and dimer formation was observed. Dimerization could be reversed by exposure to CO2 at lower temperatures. Crystallographic studies showed that, like the U complex discussed above, the O–C–O angle remained linear (179.3(5) A˚ ). In contrast to the U example, the two internal bond lengths in CO2 were much closer to one another and reflected the absence of formal reduction (O1-C1 1.162(5) and O2-C1 1.180(5) A˚ ).

1.15.3.1

Insertion into M–H

The pioneering developments of the water-gas shift reaction, coal gasification, and Fischer-Tropsch chemistry have contributed to significant ongoing interest in understanding the structure-function relationships that govern the reactivity of dihydrogen—generally via intermediate metal hydrides—with the unsaturated bonds of CO2. As a result, mechanistic studies have been pursued on the net insertion of CO2 into MdH bonds.34,46,75,76 There are two possible routes for CO2 to react with a single MdH bond: (A) normal insertion, which produces a metalloformato species, and (B) so-called “abnormal” insertion, which results in the formation of the hydroxycarbonyl M–COOH (Fig. 8). It is often informative to conceptually consider these species as analogs of an adduct of a metallocarbonyl and a hydroxide anion: they commonly have pKa values which are too high to measure and are prone to O/OH transfer reactions.45-47,77 In 1981, Darensbourg et al. reported that the anionic metal hydrides of the type [M(H)(CO)5], where M ¼ Cr, Mo, W, rapidly and quantitatively convert to the corresponding formato species (Pathway 8A) under 1 atm of CO2.78 The parallel observation of

Introduction to the Organometallic Chemistry of Carbon Dioxide

479

Fig. 8 Normal (A) and abnormal (B) insertion of CO2 into a metal hydride bond.

formate/formic acid in the electro- and photocatalytic reduction of CO2 led to interest in quantifying aspects of selectivity between H2 and formate, vide infra. In 1986, Sullivan and Meyer published a mechanistic study on the production of formate by Re(bpy) (CO)3H (bpy ¼ 2,20 -bipyridine).79 In this study, the system was determined to have an overall second-order rate law, with first-order concentration dependences on [Re(bpy)(CO)3H] and [CO2] observed across multiple solvents. The trapping agents (PPh3 and NEt3) had no effect on the observed rate law, inconsistent with a dissociative mechanism. Interestingly, an inverse kinetic isotope effect was obtained using the Re deuteride (e.g., kH/kD ¼ 0.55  0.05 in CH3CN). There were two proposed explanations: (1) a rapid pre-equilibrium followed by a rate-determining hydrogen transfer step or (2) a late transition state where CdH bond formation is nearly complete at the expense of the RedH bond. The solvent dependence showed that solvation of the dipole had an impact on initial rates, which is supportive of a highly charge-separated intermediate state. Further, substituent effects on the 4,40 -positions of the bpy ligand showed a large negative r value, suggesting positive charge buildup at the Re(I) center in the transition state. The abnormal insertion pathway (Fig. 8B) was inferred initially from observable decomposition pathways of intermediate metal hydroxycarbonyl complexes.77 Krumholz and co-workers were able to describe the true overall reaction of Fe(CO)5 decomposing to [Fe(H)(CO)4]− in the presence of a strong base, although the intermediacy of a metal hydroxycarbonyl could not be unambiguously established.80–83 Subsequent efforts to establish the mechanism of [Fe(H)(CO)4]− formation from Fe(CO)5 established indirectly that several key intermediates were likely: (1) [Fe(CO)4(CO2H)]− from hydroxide attack on the carbonyl, (2) [Fe(CO)4(CO2)]2− resulting from deprotonation of the hydroxycarbonyl, and (3) [Fe(CO)4]2− from the loss of CO2 before protonation generates the monoanionic hydride [Fe(H)(CO)4]−.84,85 Analogous reactions have been observed or implied for other hydroxycarbonyl complexes.77,86–88 Stronger evidence for abnormal insertion was reported by Schneck et al. in 2018 (Fig. 9).89 A square planar Ni(II) hydride supported by a monoanionic PNP pincer ligand ([NiH(PNP)], where PNP ¼ N(CHCHPtBu2)2), was observed to generate the formato complex [Ni(kO-O2CH)(PNP)] over the course of 14 days at room temperature, with minimal rate increases up to 10 bar of CO2 pressure. Conversely, bulk photolysis with >305 nm irradiation at room temperature under an atmospheric pressure of CO2 resulted in a 70% spectroscopic yield of the abnormal insertion product, [Ni(Z1-CO2H)(PNP)], within a few hours (quantum yield of approximately FP ¼ 5%). A Ni(II) carbonato complex made up the bulk of the mass balance of the reaction (20%) and

Fig. 9 Experimental conditions for normal vs. abnormal CO2 insertion. Reactivity of [NiH(PNP)] with CO2 under thermal and photochemical conditions, respectively, with computed free energy (DrG0(298 K), blue) and effective reaction barriers (DG{eff(298 K) red) for the respective thermal insertions. Reproduced from Schneck, F.; Ahrens, J.; Finger, M.; Stückl, A.C.; Würtele, C.; Schwarzer, D.; Schneider, S. The Elusive Abnormal CO2 Insertion Enabled by Metal-Ligand Cooperative Photochemical Selectivity Inversion. Nat. Commun. 2018, 9, 1161, with permission.

480

Introduction to the Organometallic Chemistry of Carbon Dioxide

could be obtained directly by photolysis of the isolated [Ni(Z1-CO2H)(PNP)] complex to generate [Ni(OH)(PNP)] and free CO, followed by CO2 capture to produce [Ni(kO-OCO2H)(PNP)]. Experimental and computational mechanistic studies suggested that a Ni(0) elimination product is generated by proton transfer to the N atom of the PNP framework during bulk photolysis, allowing CO2 binding in an analogous fashion to the Aresta complex discussed above. Subsequent net proton transfer generates a [Ni(Z1-CO2H)(PNP)] adduct with a computed free reaction energy of +7.3 kcal mol−1 and an activation energy of 35.6 kcal mol−1 (1 kcal mol−1 ¼ 4.184 kJ mol−1). By comparison, the metalloformato ‘normal’ insertion product which was observed over 14 days at room temperature had a computed free reaction energy change of +1.7 kcal mol−1 with an activation energy of 25.1 kcal mol−1.

1.15.3.2

Insertion into M–C

The insertion of CO2 into metal-carbon bonds (carboxylation) is a potential method for carbon chain growth in the stoichiometric or catalytic production of value-added products. In some cases, however, the discovery of novel reactions was adventitious. Work from Kohnle et al. published in 196990 showed that in studies on the dimerization of butadiene mediated by Ni/Pd/Pt-based catalysts with phosphine ligands, pressurizing the reaction vessel with CO2 led to dramatic rate enhancements in comparison to dinitrogen or argon and an enhanced selectivity for a linear 1,3,7-octatriene product. Continuing interest in controlling catalyst speciation and product distribution for reactions with unsaturated substrates like butadiene led to additional studies on the effects of CO2, during which conditions for formal insertion reaction into MdC bonds were established. In 1976, Inoue and co-workers reported an important advancement: small amounts of the lactone 2-ethylidene-hept-5-en-4-olide were observed to form when a polar solvent and a CO2 atmosphere were used under conditions where butadiene dimerization had previously been observed (Fig. 10).91 Subsequent examination of these conditions established that alternative five-membered and six-membered isomers were possible under related reaction conditions.92–94 Later work in 1983 reported by Behr et al. disclosed catalytic conditions for the synthesis of the d-lactone, 2-ethylidene-6hepten-5-olide, and the g-lactone, 2-ethyl-2,4-heptadien-4-olide.95–97 Interestingly, the d-lactone isomerized to the g-lactone under the reaction conditions, but more slowly than the initial cyclization reaction. Later work by Braunstein and co-workers in 1988 presented a more exhaustive study of phosphine ligand effects and product distribution.98 Concurrent work by Jolly and co-workers focused on establishing key catalytic intermediates, confirming the relevance of both a PddC bond and the insertion of CO2 into that bond.99–101 Ukai et al. reported a Rh-catalyzed carboxylation reaction in 2006, which relied on a transmetallation scheme using boronic acids and CO2 to generate aryl- and alkenyl-based carboxylate products.102 With mild heating in the presence of a base and a boronic acid, a transient RhdC bond was generated from BdC bond activation. This transient RhdC bond could subsequently undergo CO2 insertion to produce an intermediate Rh-bound carboxylate species. An additional equivalent of the boronic acid produced the boronic ester product of the carboxylated alkenyl and aryl substrates, completing the catalytic cycle. Given the widespread availability of boronic acid substrates, selective carboxylation reactions directly from CO2 offer convenient routes to carboxylic acids.103 Subsequent development of the transmetallation strategy by Yeung and Dong reported the use of Aresta’s complex as a catalyst for reductive CO2 coupling with aryl and alkylzinc substrates.104 This is a conceptual extension of the Negishi cross-coupling reaction105 with CO2 as a co-substrate and leverages Aresta’s complex to overcome the thermodynamic penalty of CO2 activation.

Fig. 10 Minor lactone side-product observed by Inoue and co-workers,91 suggesting that insertion of CO2 into a PddC bond had occurred, with subsequent thermal isomerization and cyclization.

Introduction to the Organometallic Chemistry of Carbon Dioxide

481

Fig. 11 Plausible reaction mechanism for the carboxylation of alkyl and alkyl organozinc precursors catalyzed by Ni-phosphine species.104 A key proposed intermediate is the isolable Ni(PCy3)2(Z2-CO2). Note that other triaryl and trialkylphosphines give viable catalyst systems as well.

Not only did the method show excellent selectivity under mild conditions at 0  C and atmospheric CO2 pressure, but it was also functional with either Ni or Pd as the catalytic metal center. In the mechanism, a M(0) species (where M ¼ Ni or Pd) was proposed to coordinate CO2 in Z2-CO2 fashion, formally reducing it by two electrons (Fig. 11).104 Following transmetallation, reductive elimination would generate the carboxylic acid derivatives as Zn salts, regenerating the Ni(0) starting species. Later optimization with salt additives showed increased functional group tolerance and suitability of Cu complexes and transition metal-free conditions for these types of transmetallation reactions.106–109 An expanded discussion of additional examples of catalytic carboxylation reactions is found in later sections.

1.15.3.3

Insertion into M–N

Complete characterization of the insertion of CO2 into MdN bonds was first reported by Chisholm and Extine in 1974.110 Exposing a hydrocarbon solution of hexakis(dimethylaminato)tungsten, W(NMe2)6, to an excess of CO2 at room temperature resulted in quantitative conversion to the corresponding tris(N,N-dimethylcarbamato) species, W(NMe2)3(O2CNMe2)3. The presumed intermediate mono- and bis(N,N-dimethylcarbamato) complexes were detected spectroscopically, although their conversion to the tris species is rapid. Cryoscopic mass spectra using samples prepared using isotopically enriched CO2 showed the uptake of three equivalents of CO2; 1H NMR spectra obtained on samples with normal and isotopically enriched 13CO2 enabled the peak assignments from the observed 3J13C-1H coupling constants. The observation of a single IR absorption band at 1632 cm−1 led to a speculative assignment of a fac-configuration, which was validated through X-ray crystallography (Fig. 12). Despite the steric congestion in W(NMe2)6, its reaction with CO2 is rapid. Interestingly, exposing W(NMe2)3(O12 2 CNMe2)3 to 13 CO2 results in a distribution of isotopologues, indicative of reversible interconversion of the carbamates with CO2. The authors proposed that the nucleophilicity of the N lone pair of the dimethylamido groups was the limiting factor for CO2 uptake. For the parent species, W(NMe2)6, the ability of all six dimethylamido equivalents to engage in a p-bonding interaction with the d0 W center is limited by the preference of the metal center to accommodate only 18 electrons. Relative to dimethylamido, N, N-dimethylcarbamato is anticipated to be a weaker p donor, allowing better p overlap for dimethylamido with each CO2 insertion.

Fig. 12 Crystallographic structure of W(NMe2)3(O2CNMe2)3, CCDC 1265058. H atom positions not determined; key bonding metrics: W–O1 2.041(6) A˚ , W–N2 1.922(7) A˚ , O1–C1 1.30(1) A˚ , and C1–O2 1.24(1) A˚ ; O1–C1–O2 122.8(9) .

482

Introduction to the Organometallic Chemistry of Carbon Dioxide

This reaches a limit at three equivalents, where the p-bonding from N to W is maximized, resulting in a minimization of the WdN bond lengths (1.922(7) A˚ ), relative to increased WdO bond lengths (2.041(6) A˚ ), which have a comparatively weaker p bond (Fig. 12). Reports have been disclosed where comparable reactivity was observed resulting from the net insertion of CO2 into nitrogen-metal bonds for Ti, Zr, V, Nb, Ta, Sn, Si, Ge, P, As, Sb, Zn, Al, Fe, Ce, U, Mo, W, and Mg complexes.111–138 Interested readers are also directed to general reviews of the related chemistry of metal carbamates.44,139,140

1.15.3.4

Insertion into M–O

In the presence of hydroxide CO2 is readily captured and converted into carbonate. Many CO2 “scrubber” technologies make use of this process with alkali and alkaline earth hydroxides and carbonates used to facilitate a complete capture and recovery cycle.141–144A related “fixation” process occurs in mammalian tissues, mediated by the Zn-based carbonic anhydrase metalloenzyme.145–147 This enzyme catalytically converts CO2 to carbonate and back again, enabling net CO2 transport to the lungs as a part of the respiratory cycle, with a Zn(II) hydroxide species as a key intermediate. Net CO2 insertion to generate functionalized carbonate moieties within transition metal complexes is known to occur for a variety of M–OR motifs, where R ¼ H, alkyl, or aryl.46 The convergence of mechanistic studies and the related interest in replicating the function of the carbonic anhydrase enzyme has lead to synthetic variants containing Zn–OH moieties capable of reversible CO2 capture based on this reaction. In a series of papers in the early 1990s, Parkin and co-workers reported a viable carbonic anhydrase mimic based on a monoanionic tris-pyrazolylborate (pz) ligand framework (Fig. 13).148–150 For the sterically demanding structural model, (Z3-HB(3-tBu-5-Mepz)3)ZnOH, CO2 capture was found to be highly reversible. Line broadening observed in 1H NMR studies under CO2 atmosphere was used to estimate a rate constant on the order of 102–103 M−1 s−1 for the net insertion of CO2 into the ZndOH bond to generate (Z3-HB(3-tBu-5-Mepz)3)Zn(Z1-OCO2H). Complete characterization of this CO2 capture equilibrium in this initial study was difficult owing to the dynamic behavior of the equilibrium, as well as the existence of an additional species generated by a dehydration reaction involving the starting material (Z3-HB(3-tBu-5-Mepz)3)ZnOH and the CO2 insertion product (Z3-HB(3-tBu-5-Mepz)3)Zn(Z1-OCO2H) to produce one equivalent of H2O and a dimeric species, ((Z3-HB(3-tBu-5-Mepz)3)Zn)2(m2-Z2-O3C). Various derivatives of the tris-pyrazolyl framework have been used as ligand platforms for additional Zn-based enzymes,151 as well as other metalloenzymes.152–154 Additional model ligand platforms for the carbonic anhydrase enzyme were developed with macrocyclic polyamine ligands, beginning with a report in 1990 by Kimura et al.155–157 In initial studies, reactions catalyzed by the native carbonic anhydrase enzyme were used to evaluate structural and activity comparisons with the model complexes, instead of CO2 to bicarbonate conversion. Subsequent mechanistic studies in 1993 on CO2 capture by the cyclic triamine-based model complex [(1,5,9-triazacyclododecane)Zn(OH)][ClO4] established kinetic and thermodynamic parameters for CO2 conversion to bicarbonate, as well as reaction parameters for bicarbonate dehydration.158 Zhang and Eldik continued these efforts in 1995 with a tetraaza-macrocyclic variant, [Zn(cyclen)(OH2)][ClO4]2, where cyclen ¼ 1,4,7,10-tetraazacyclododecane. To date, the [Zn(cyclen) (OH2)][ClO4]2 is the most active synthetic Zn-based catalyst known for catalytic conversion of CO2 to bicarbonate, with a second-order catalytic rate constant, kcat, of 3.3  0.1  103 M−1 s−1. The relative synthetic ease of producing the cyclen ligand, the high rate of catalytic activity of the corresponding Zn complex, and the overall robustness of this platform has even enabled testing of catalytic CO2 capture under industrially relevant conditions (Fig. 14).159 This propensity of metal alkoxide complexes to undergo facile CO2 insertion has been harnessed in the context of catalytic conversions. In 2013, Ishitani reported favorable CO2 capture (Keq ¼ 1.7  103 M−1) by a Re(bpy)(CO)3(OR) complex, where RO− is monodeprotonated triethanolamine.160 It was subsequently established that this capture mechanism occurred even at low CO2 concentrations, for both Re and the Mn complexes with different 4,40 -substituted bpy ligands, allowing reaction parameters to be compared across multiple derivatives.161 In a culmination of these efforts to concentrate dilute gas streams for chemical

Fig. 13 Comparison between the structure of the active site of Zn-based carbonic anhydrase metalloenzyme (left) and the structure of the model complex Zn(Z3-HB(3,5-RR’pz)3)OH (right), R ¼ 3-tBu-5-Me. Adapted with permission from ref Looney, A.; Han, R.; McNeill, K.; Parkin, G. Tris(Pyrazolyl)Hydroboratozinc Hydroxide Complexes as Functional Models for Carbonic Anhydrase: On the Nature of the Bicarbonate Intermediate. J. Am. Chem. Soc. 1993, 115, 4690–4697. Copyright 1993 American Chemical Society.

Introduction to the Organometallic Chemistry of Carbon Dioxide

483

Fig. 14 Overall capture scheme catalyzed by [Zn(cyclen)(OH2)][ClO4]2 in basic aqueous solution. Adapted with permission from Floyd, W.C.; Baker, S.E.; Valdez, C.A.; Stolaroff, J.K.; Bearinger, J.P.; Satcher, J.H.; Aines, R.D. Evaluation of a Carbonic Anhydrase Mimic for Industrial Carbon Capture. Environ. Sci. Technol. 2013, 47, 10049–10055. Copyright 2013 American Chemical Society.

conversion, electrocatalytic reduction of CO2 could be performed on gas mixtures containing as little as 1% CO2 in argon while maintaining excellent selectivity for CO (the two-electron/two-proton reduction product of CO2; H2O formed as co-product) over a 24 h period.162

1.15.3.5

Metathesis

In spite of the relatively short carbon-oxygen bond distances in CO2, the bonds are weak relative to other unsaturated bonds, e.g., ethylene. As a result, despite the kinetic inertness of CO2, the molecule is poised to undergo multiple bond metathesis reactions with other unsaturated small molecules. The first example was reported in 1963 by Bigorgne and Rassat,163 who disclosed the conversion of a zero-valent mononuclear nickel tetra(arylisocyanide) complex, Ni(CNAr)4 (Ar ¼ 2,6-Me2-C6H3) in the presence of carbon dioxide and Li+ in THF solution to Ni(CO)2(CNAr)2 and two equivalents of the corresponding aryl isocyanate (O]C] NAr). In 1998, Kubiak and co-workers conducted a mechanistic study on this reaction, using 13C isotopologues of the isocyanide ligands and CO2 substrate.164 With isotopically enriched CO2 present, the Ni(CO)2(CNAr)2 product shows no difference in the observed IR frequencies, but the aryl isocyanate frequencies shift by 23 cm−1 to lower energy. A second labeling study with the methylisocyanide isotopologue Ni(13CNMe)4 generated Ni(13CO)2(13CNMe)2 as the exclusive product; no incorporation of 13C was observed in the methyl isocyanate co-product. These studies also established that the reaction was catalytic with respect to Li+, as little as 1% in solution could mediate the reaction, and the use of 12-crown-4 to bind Li+ resulted in a significantly diminished rate of reaction. The results of these kinetic and mechanistic studies led the authors to propose that Li+ ions served to stabilize the development of negative charge on the CO2 oxygen atoms, reminiscent of the Co(salen) chemistry described above60 (although the authors cautioned that they had not observed evidence of this directly). The overall reaction mechanism was proposed to rely on the carbenoid resonance form of the isocyanide, resulting in significant lone pair character on the N atom. This allows the N atom of the isocyanide to engage in a nucleophilic attack of the CO2, initiating an exchange that pairs the NR (R ¼ Ar or Me) fragment of the isocyanide with a CO fragment from CO2, generating the metal carbonyl (Fig. 15).

Fig. 15 Reaction scheme of the initial Li+-catalyzed metathesis reaction between a Ni tetrakis(isocyanide) complex and CO2. This cycle repeats with a second equivalent of CO2 and the Ni(CO)(CNR)3 intermediate to produce the final Ni(CO)2(CNR)2 product, where R ¼ 2,6-Me2-C6H3 or Me.164

484

Introduction to the Organometallic Chemistry of Carbon Dioxide

Fig. 16 Reaction scheme depicting atmospheric and high-pressure pathways for the dinuclear Ni complex Ni2(m2-CNMe)(CNMe)2(dppm)2, and CO2.166

The assignment of the carbenoid isomer of the metal isocyanide as a key intermediate was based on earlier work reported by Kubiak and co-workers with the dinuclear Ni-based cradle complex, Ni2(m2-CNMe)(CNMe)2(dppm)2 [dppm ¼ bis (diphenylphosphino)methane].165–167 For this dinuclear complex, crystallographic studies revealed that the bridging isocyanide ligand adopts a bent CdNdCMe bond angle of 129.9(6) , suggesting a high degree of activation and significant carbenoid character.168 Indeed, the N atom demonstrated significant Lewis basicity, comparable to that of ammonia under similar conditions. Although the complex is unreactive toward CO2 at 1 atm, the addition of catalytic amounts of NaPF6 (1%) resulted in the formation of a neutral carbamate-like adduct, Ni2(m2-CN(Me)CO2)(CNMe)2(dppm)2 (Fig. 16). Under elevated pressures of CO2 (ca. 100–150 atm) at room temperature over 48 h, the dinuclear tricarbonyl complex Ni2(m2-CO)(CO)2(dppm)2 is generated, with a presumed carbodiimide polymer comprising the mass balance of the reaction. At shorter reaction times under these conditions, both Ni2(m2-CN(Me)CO2)(CNMe)2(dppm)2 and Ni2(m2-CO)(CO)2(dppm)2 are observed, suggesting that the former is an intermediate during the reaction. The authors proposed that this made methylisocyanate a probable intermediate to the carbodiimide polymer, as the CO2 adduct was converted to the tricarbonyl complex. Isotopic labeling studies of the same reaction with Ni2(m2-CN(Me)CO2)(CNMe)2(dppm)2 at elevated CO2 pressures confirmed that the source of the C atom in the tricarbonyl product is the isocyanide ligand and not CO2, analogous to the metathesis reaction described in Fig. 15.165,166 These observations are consistent with the proposed methyl isocyanate intermediate. Interestingly, photolysis of THF solutions of Ni2(m2-CNMe)(CNMe)2(dppm)2 with 355 nm light under 1 atm CO2 was also found to produce Ni2(m2-CN(Me)CO2)(CNMe)2(dppm)2 with an overall quantum yield of F355 ¼ 0.05.167 This wavelength corresponded to the lowest energy electronic absorption of the complex and enabled transient absorption studies. These studies identified that upon irradiation at 355 nm, charge transfer from the electron-rich Ni-Ni manifold into the p antibonding orbitals of the m2-CNMe ligand occurred to generate a singlet excited state [MLCT]1, followed by intersystem crossing to generate a relatively long-lived triplet [MLCT]3. The lifetime of the triplet [MLCT]3 product (300 ms) was sufficient to enable reactivity with CO2, forming the same Ni2(m2-CN(Me)CO2)(CNMe)2(dppm)2 species observed at elevated pressures and implying a relative increase in nucleophilic character at the N atom in the excited triplet state. In 1996, Sita et al. reported a related metathesis reaction involving Ge and Sn bis(amide) complexes wherein CO2 was converted to trimethylsilyl isocyanate and carbodiimide products under elevated pressure, with co-generation of the corresponding siloxide complexes.169 Subsequent mechanistic studies suggested that the overall mechanism required generation of a more reactive monomeric tin(II) species prior to CO2 metathesis, implying that the coordination of multiple bridging alkoxides has an inhibitory effect.170 Ghosh and Samuelson subsequently reported that titanium isopropoxide could mediate a related metathesis reaction, using carbodimides and CO2 as substrates to generate isocyanate and carbamate products.171 Analogous metathesis to produce isocyanates and diimides have also been reported for titanium imides exposed to CO2.172 More recently, Mokhtarzadeh et al. disclosed evidence of a metal-mediated variant of the isocyanide reaction above with a comparable dianionic Fe complex.173 Beginning from the sterically encumbered tetraisocyanide dianion [Na2][Fe(CNArMes2)4], where ArMes2 ¼ 2,6-(2,4,6-Me3C6H2)2C6H3, exposure to a CO2 atmosphere results in the formation of the monocarbonyl complex Fe(CO)(CNArMes2)4. Note that the tetrahedral iron(-II) starting material is isoelectronic with the neutral Ni-based tetraisocyanides discussed above. The change in charge and final product formulation implied that CO2 coordination to the Fe center was occurring prior to a reductive disproportionation reaction, in which a second equivalent of CO2 was converted to carbonate, as the Fe-bound equivalent was reduced to CO. In an attempt to intercept the Fe-bound CO2 species, the sterically hindered silylating reagent Me3SiOTf (which displayed no intrinsic activity toward the dianion under these conditions, OTf ¼ triflate) was added to the reaction mixture. The isolated product from this reaction was a contact ion pair containing a terminal Fe carbyne, Na[Fe^CN((ArMes2) (C(O)OSiMe3))(CNArMes2)3]. Decarboxylation could be achieved inefficiently with the addition of protic reagents or reversed in a stoichiometric way with two equivalents of Na[HBEt3], resulting in the regeneration of the starting material with formate and HSiMe3 as co-products.173 The metal-assisted mechanism proposed by Mokhtarzadeh has also been proposed by Broere et al. based on kinetic evidence and THF solvent inhibition in the reaction of low-coordinate Fe silylamide complexes with CO2 to give trimethylsilyl isocyanate and bridging Fe siloxide compounds as products.111

Introduction to the Organometallic Chemistry of Carbon Dioxide

1.15.4

485

Homogeneous catalysis with CO2

The most common products from the direct reduction of CO2 by molecular catalysts are formic acid (HCO2H) and CO, net 2H+/2e− transformations.10,174 Formic acid can be used for the reversible storage of molecular hydrogen,175,176 as an organic hydride reagent,177,178 and in fuel cells.10 CO has many industrial applications, including the Fischer-Tropsch production of liquid fuels,179–186 hydroformylation reactions,187–189 and acetic acid production.190 The direct reduction of CO2 can instead utilize H2 as the source of reducing equivalents.27,191,192 Notably, formate/formic acid products from CO2 hydrogenation are often viable intermediates to more highly reduced molecules like methanol under reaction conditions; carbonates and carbamates are also viable substrates under these conditions.139 Currently, C2 products from CO2 can be obtained electrochemically only by using heterogeneous Cu-based materials as catalysts, albeit with less-than-quantitative selectivity.24,193–198 Other electrocatalysts generally produce CO but lack an efficient pathway to CdC bond formation. In principle, this CO could be used in the Fischer-Tropsch reaction, combining it with H2 to form hydrocarbons.179-186,199 Although CO2 can be used industrially to make CO via variants of the Boudouard200 reaction at high temperatures (ca. 800  C), this is not the primary source of the CO used in Fischer-Tropsch-based reactions. Heterogeneous electrocatalysts that show excellent selectivity for CO are known and are a feasible source for this CO,201 if renewable electricity costs continue to decrease or carbon emissions are taxed at higher rates.202 As a co-substrate in catalytic reactions, the CO2 molecule can be incorporated in its entirety, generating products like carbamates, carbonates, and carboxylates.35,203 Catalytic carboxylation reactions are known to occur with a variety of metal centers under relatively mild conditions.28,103,204–213 Abundant examples also exist where the overall co-substrate incorporation process generates polymeric materials in lieu of monomeric ones.40,214–217 Comprehensive reviews of catalytic CO2 conversions are cited throughout this section, and selected examples will be discussed in greater detail below. Heterogeneous catalysts are beyond the scope of this article but are a related area of significant interest.22,23,29,218

1.15.4.1

Hydrogenation

In 1976, Inoue et al. reported the first example of homogeneous hydrogenation of CO2 to formic acid under catalytic conditions.219 Under elevated pressures of CO2 (25 atm) and H2 (25 atm) with triethylamine present as a base, they found that Pd(diphos)2 [diphos ¼ bis(diphenylphosphino)ethane] could generate formic acid with a turnover frequency (TOF) of 3.5 h−1 at 140  C. It should be emphasized that TOF is a measure of kinetic activity and does not offer insight on catalyst lifetime or stability. H2Ru(PPh3)4 was also a viable catalyst for formic acid under these conditions, showing a slightly higher TOF of 4.4 h−1. Bases are common additives for the most catalytically active systems, improving the enthalpy of the reaction by favoring the formation of formate salts instead of formic acid from H2 and CO2 alone, which is thermodynamically unfavored.220 Shortly after this report, Kudo et al. disclosed that a PdCl2 catalyst that could hydrogenate potassium bicarbonate to potassium formate.221 Subsequent optimization and kinetic analysis of potassium formate production by PdCl2 under aqueous alkaline conditions suggested that the bicarbonate anion was a key intermediate in the overall transformation.222 Since these initial reports, the activity and catalyst stability for CO2 hydrogenation have rapidly improved with the discovery of additional catalyst platforms.27,223 Given the intrinsic limitations of working with gaseous substrates and the propensity of formic acid to decompose under catalytic conditions, significant effort has focused on systems that can operate in supercritical CO2 and/or in biphasic reaction systems.224 An increase in turnover number (TON) to 3400 was achieved by Gassner and Leitner in 1993 using a catalyst system generated in situ by combining [Rh(cod)Cl]2 (cod ¼ 1,4-cyclooctadiene) with added tppts (tppts ¼ trisodium salt of triphenylphosphine-3,30 ,300 -trisulfonic acid).225 Reported TONs generally reflect the lifetime of the catalyst system before complete loss of activity. Noyori and co-workers were the first to disclose a system with supercritical CO2 as the reaction medium, where a TON of 7200 was achieved for formic acid.226 With RuCl2(PMe3)4 as the catalyst, triethylamine as the base, and a supercritical mixture of H2 and CO2 at 50  C, an initial TOF of 1400 h−1 was achieved. Further optimization of the general catalyst system with reactions coupled to formic acid, enabled generation of alkyl formates and formamides under these conditions.227 More recently, Tanaka et al. described a groundbreaking Ir(III)-pincer system where a TON of 3,500,000 and a TOF of 150,000 h−1 could be achieved in the production of potassium formate using H2 and CO2 in THF, with potassium hydroxide as a base.228 The proposed catalytic cycle invoked a dearomatization of the pyridine ring in the ligand backbone, prior to the regeneration of the trihydride complex (Fig. 17). Schmeier et al. subsequently reported an alternative way for tridentate ligands to participate in an Ir-catalyzed CO2 hydrogenation cycle by capitalizing on hydrogen-bonding interactions with an –NH– group in the ligand backbone.229 Jessop and co-workers applied a high-throughput screening methodology that identified the first highly active CO2 hydrogenation catalyst outside the platinum group.230 Both FeCl3 and NiCl2 became active for CO2 hydrogenation with added Cy2PCH2CH2PCy2 as a ligand. Following this advancement, Federsel et al. reported a well-defined [Fe(P(CH2CH2PPh2)3)] [BF4]2 catalyst that was capable of reducing bicarbonates and carbon dioxide to formates, alkyl formates and formamides, depending on reaction conditions.231 High-pressure in situ NMR studies supported a mechanistic proposal where heterolytic H2 cleavage occurred to form an Fe(II) hydride. Interestingly, these studies also established that an Fe-bound CO2 intermediate preceded the insertion reaction and that an off-cycle equilibrium exists with [FeH(H2)(P(CH2CH2PPh2)3)]+ at the initiation point of the catalytic cycle (Fig. 18). Milstein and co-workers reported another example of Fe-catalyzed CO2 hydrogenation that relied on a PNP pincer ligand analogous to the Ir system discussed above.232

486

Introduction to the Organometallic Chemistry of Carbon Dioxide

Fig. 17 Plausible mechanism for hydrogenation of CO2 using an iridium trihydride complex. Adapted with permission from Tanaka, R.; Yamashita, M.; Nozaki, K. Catalytic Hydrogenation of Carbon Dioxide Using Ir(III)−Pincer Complexes. J. Am. Chem. Soc. 2009, 131, 14168–14169. Copyright 2009 American Chemical Society.

Fig. 18 Proposed catalytic cycle for CO2 hydrogenation for a well-defined [Fe(P(CH2CH2PPh2)3)][BF4]2 catalyst based on high-pressure NMR studies in THF.231

Hydrogenation of CO2 to the more reduced product, formaldehyde, was first disclosed in 1989 by Taqui Khan et al., who found that K[Ru(EDTA-H)Cl]•2H2O, where EDTA-H is the monoprotonated form of ethylenediaminetetraacetic acid, was an active catalyst.233 The maximum rates of formation for formic acid and formaldehyde were achieved at elevated pressures with a 1:1 mixture of CO2 and H2 at 40  C. Decomposition of both products to CO was appreciable under reaction conditions and the authors proposed that the mechanism included an abnormal insertion of CO2 into the MdH bond, although direct evidence of this pathway was not obtained. It is also worth noting that disproportionation of formic acid to formaldehyde, water, and CO2 is thermodynamically feasible at room temperature and could be the source of the observed formaldehyde.234 Miller et al. were able to capitalize on the thermodynamic favorability of formaldehyde disproportionation using an Ir-based catalyst system, achieving methanol as a final product with formaldehyde as a discrete intermediate (Fig. 19).234 Aqueous solutions of formic acid containing

Fig. 19 Proposed catalytic cycle234 for the disproportionation of three equivalents of formic acid to methanol in aqueous solutions by [Cp Ir(bpy)(H2O)] [OTf]2 (Cp ]pentamethylcyclopentadienyl, bpy]2,20 -bipyridine). Adapted with permission from Wang, W.-H.; Himeda, Y.; Muckerman, J. T.; Manbeck, G. F.; Fujita, E. CO2 Hydrogenation to Formate and Methanol as an Alternative to Photo- and Electrochemical CO2 Reduction. Chem. Rev. 2015, 115, 12936–12973.

Introduction to the Organometallic Chemistry of Carbon Dioxide

487

[Cp Ir(bpy)(H2O)][OTf]2 (Cp ¼pentamethylcyclopentadienyl, bpy ¼ 2,20 -bipyridine) were shown to produce methanol upon heating. The reaction sequence was proposed to start with the dehydrogenation of formic acid to produce an intermediate Ir hydride, which could reduce formic acid to formaldehyde as an intermediate, with methanol production occurring upon the dehydrogenation of a second equivalent of formic acid. There is continued parallel interest in the hydrogenation of CO2 to methanol, a reaction that is thermodynamically favorable, unlike formic acid production which is unfavorable unless deprotonated to formate. Milstein and co-workers reported the first molecular system for the related transformation of carbonates, carbamates and formates to methanol.235 The key to this transformation was the use of a Ru(II) PNN pincer complex that was partially dearomatized,236 which could activate H2 to initiate the catalytic cycle and was proposed to participate in hydrogen transfer during substrate reduction: three distinct cycles of dearomatization following rearomatization by H2 were proposed to occur (Fig. 20). Han et al. developed an alternative system for the conversion of ethylene carbonate to 1,2-ethylene glycol and methanol, a substrate which is more readily generated from CO2 than dimethylcarbonate.237

1.15.4.2

Electrochemical reduction

The most common products of electrochemical CO2 reduction by molecular species are CO and formate/formic acid. These reduction products are generally achieved through separate activation pathways. The primary path to CO is nucleophilic attack by a reduced metal center on the electrophilic C atom of CO2 in the presence of a suitable proton donor results in the formation of an intermediate hydroxycarbonyl (Fig. 21, i), from which CdOH bond cleavage can occur to produce a metal carbonyl (ii). Note that the abnormal insertion pathway could also be possible, but has not been directly verified in electrocatalysis.89 For formate/ formic acid the formation of an intermediate metal hydride species (pKa(1), iii) occurs first, after which hydride transfer to the C atom of CO2 (iv) can occur as part of a net insertion reaction. For the CO-producing mechanism to be viable, selectivity for

Fig. 20 Proposed catalytic cycle236 for the reduction of dimethylcarbonate to methanol by a dearomatized Ru PNN pincer complex under H2 atmosphere. Adapted with permission from Wang, W.-H.; Himeda, Y.; Muckerman, J. T.; Manbeck, G. F.; Fujita, E. CO2 Hydrogenation to Formate and Methanol as an Alternative to Photoand Electrochemical CO2 Reduction. Chem. Rev. 2015, 115, 12936–12973. Copyright 2015 American Chemical Society.

Fig. 21 Mechanisms for CO and HCO−2 production by mononuclear active sites.

488

Introduction to the Organometallic Chemistry of Carbon Dioxide

CO2 over H+ is required. For the formate-producing mechanism, the pKa of the intermediate metal hydride (pKa(2), Fig. 21) must be lower than the pKa of the acid to inhibit competing H2 formation and the hydricity of the metal hydride should thermodynamically favor net CO2 insertion.238 Alternatively, if no suitable proton donor is present, a second equivalent of CO2 can act as an oxo anion acceptor, generating carbonate as a co-product to CO in an overall reductive disproportionation of two equivalents of CO2 when two electrons are provided. Selected discoveries and advances in electrochemically driven molecular catalysis for these reactions follow below in approximate chronological order, and photocatalytic behavior has been summarized in detail elsewhere.20,26,239–243 The first report of electrochemical CO2 reduction with a discrete molecular complex was reported in 1974 by Mehitsuka et al. who deposited metal phthalocyanines onto graphite electrodes for characterization (Fig. 22).260 In these studies, the Co and Ni derivatives showed the highest activity. Although not a homogeneous system, this represented the first example of electrochemical conversion where the active material was not a metal or a mercury-based electrode (hanging drop or amalgam).261–270 The first example of homogeneous catalysis for CO2 reduction came in 1977 from Kazuya et al. using tetrasulfonated phthalocyanines in an aqueous solvent system (again Co and Ni showed the greatest activity).271 Takahashi et al. subsequently extended the characterization of homogeneous catalytic behavior to Co porphyrins using tetracarboxylated and tetrasulfonated derivatives in 1979.272 Interestingly, this study found that the Fe derivatives were inactive in Clark-Lubs buffer from pH 8.0 to 10.3. However, product efficiencies were not reported in either case.271,272 Fisher and Eisenberg then reported in 1980 that Ni and Co complexes with cyclam-based ligands—as well as a macrocyclic pyridyl diimine example for Ni—were active catalysts for CO2 reduction (cyclam ¼ 1,4,8,11-tetraazacyclotetradecane).244 In a mixture of water and acetonitrile, a [Ni(cyclam)]2+ derivative showed an overall 98% Faradaic efficiency for a 2:1 product mixture of CO and H2. Beley et al. improved on this result with a purely aqueous solvent system with a mercury electrode, achieving 99% efficiency in 1984.273 The first example of near quantitative selectivity by a molecular species came in 1982, when Tezuka et al. found that a iron-sulfur clusters [Fe4S4(SR)4]2− (R ¼ CH2C6H5 or C6H5) could catalyze CO2 reduction in dry DMF, producing formate with as high as 93% Faradaic efficiency.245 The authors proposed that the tetraalkylammonium electrolyte salts degraded to provide the necessary proton and noted the overall reaction efficiency for formate diminished when water was added.274 In 1984, Hawecker et al. reported that Re(bpy)(CO)3Cl was an active and selective electrocatalyst for CO production (98% Faradaic efficiency) during electrochemical CO2 reduction in a DMF-H2O (10%) solvent mixture.246 Notably, the authors had reported that this Re complex was also an active single-component photocatalyst the previous year.275 A Rh-based electrocatalyst for CO2 reduction with bis(diphosphine) ligands was also reported in 1984 by Slater and Wagenknecht, but achieved only 42.5% efficiency for formate under optimized conditions.276 Ishida et al. disclosed a [Ru(bpy)2(CO)2]2+ electrocatalyst for CO2 reduction in 1985, which showed selectivity primarily for CO in 1:1 mixtures of DMF and H2O under acidic conditions, but with 1:9 mixtures of DMF and H2O under more basic conditions produced near equimolar amounts of CO, H2, and formic acid.247 Further optimization by Ishida et al. in 1987 improved the selectivity for formate to 84.3% by systematically varying the proton donor.277 Meyer and co-workers reported that the comparable osmium-based complex was also active for CO2 reduction in 1992, producing primarily CO under anhydrous conditions in CH3CN and up to 22% formate as a co-product when H2O was added.278 In 1988, the first report of Fe tetraarylporphyrins as active catalysts for CO2 reduction in dry DMF to CO appeared from Savéant and co-workers.279 Reports improving the reaction activity and selectivity for CO through the inclusion of Mg2+ cations and Brønsted acids followed in 1991248 and 1996,249 respectively, from the same research group. Recently, significant advancements in improving catalytic performance have been achieved through synthetic modification of the porphyrin ligand framework with pendent functional groups.280–285 Ishida et al. disclosed an additional noteworthy advancement in 1990, reporting the activity of two types of [Ru(bpy) (CO)2(Cl)2] complexes for CO2 reduction, producing mixtures of CO and formate as the primary products.250 Later work showed that under reducing conditions, this species was prone to electropolymerization, forming a [Ru(bpy)(CO)2]n metallopolymer286 which remained active for electrocatalytic CO2 reduction (Fig. 23).287 Variants of this catalyst are often used as catalytic components of artificial photosynthetic systems, where the metallopolymerization reaction is less of a concern.26 The complex can function as a discrete molecular catalyst, however, if sufficient steric bulk is included on the ligand framework, inhibiting polymerization.86 The catalytic properties of the Os congener of the metallopolymer for CO2 reduction (CO major product, formate minor product) were reported in 2001.288 DuBois and co-workers disclosed in 1991 Pd-based catalysts with triphosphine ligands, with efficiencies for CO as high as 85% over 130 turnovers in DMF solution with HBF4 present as an acid.251 In 1994, Nagao et al. reported the first example of a [Ru(bpy) (tpy)CO]2+ electrocatalyst for CO2 reduction (Fig. 24).252 In CO2 saturated ethanol/water mixtures a mixture of CO, formic acid, and H2 were produced. Interestingly, following low temperature (−20  C) electrolysis, four- and six-electron reduction products (methanol, acetaldehyde, and glycolic acid) were reported, although these results have not been replicated. Subsequent iterations of the design and reaction conditions have resulted in significant improvements in activity and selectivity by substituting pyridines with carbene motifs.289 Substitution within the bpy ligand has important consequences with respect to geometric isomers, as the carbene fragment can be placed trans to the site of CO2 binding, which increases activity but has a minimal effect on the Faradaic efficiency for CO.290 Interestingly, the Fe congener also demonstrates similar activity and selectivity (>95%) for CO2 reduction to CO at low overpotentials, although at more modest rates.291 Interestingly, the Fe congener also undergoes isomerization reactions involving the NHC-pyridine ligand, as was observed for Ru.292 A related covalently linked pentadentate polypyridine Fe variant with an alternative geometric configuration demonstrated significantly higher activity for CO2 reduction to CO.293 Self-exchange

Introduction to the Organometallic Chemistry of Carbon Dioxide

Fig. 22 Notable catalyst advancements for electrochemical reduction of CO2. Catalytic properties for homogeneous compounds summarized in Table 1.

489

490

Introduction to the Organometallic Chemistry of Carbon Dioxide

Table 1

Summary of homogeneous electrocatalytic properties for complexes displayed in Fig. 22 according to initial report.

Complex

Operating potential (V)

Major product

Faradaic efficiency (%)

Conditions

References

2 3 4 5-CH2C6H5 5-C6H5 6 7 7 8 8 9 10-R 11

−1.6 V vs. SCE −1.6 V vs. SCE −1.3 V vs. SCE −2.1 vs. SCE −2.0 vs. SCE −1.25 vs. NHE −1.50 vs. SCE −1.50 vs. SCE −1.70 vs. SCE −1.70 vs. SCE −1.30 vs. SCE −1.38 vs. Fc+/Fc −1.60 vs. Ag/AgNO3

65 47 44 81 93 98 48 29/39 75 96 88 85 35/30

0.1 M LiClO4 in H2O/CH3CN 2:1 v/v or H2O only 0.1 M KNO3 in H2O/CH3CN 2:1 v/v or H2O only 0.1 M KNO3 in H2O/CH3CN 2:1 v/v 0.1 M NBu4BF4 in DMF 0.1 M NBu4BF4 in DMF 0.1 M NBu4ClO4 in DMF/H2O (10% v/v) pH 6.0 H2O NaOH-H3PO4 H2O/DMF 9:1 v/v; pH 9.5 NaOH-H3PO4 15 mM Mg2+ (Mg anode); 0.1 M NEt4ClO4 in DMF 1.37 M CF3CH2OH; 0.1 M NEt4ClO4 in DMF MeCN/water 4:1 with 0.1 M NBu4ClO4 0.1 M HBF4 in DMF with 0.3 M NEt4BF4 DMF/H2O 2:8 v/v; pH 9 with LiCl

244 244 244 245 245 246 247 247 248 249 250 251 252

12 13 14 14 15-W 16 17 18

−1.70 vs. Ag/AgNO3 −1.45 vs. NHE −1.20 vs. SCE −1.20 vs. SCE −2.30 vs. SCE −1.65 vs. SCE −2.15 vs. Fc+/Fc −2.10 vs. Fc+/Fc

CO CO CO HCOO− HCOO− CO CO CO/HCOO− CO CO CO CO CO/ HCOO− CO HCOO− HCOO− HCOO− CO CO HCOOH CO

100 85 96 94 109 96 99 96

MeCN/H2O 95:5 with 0.1 M NBu4ClO4 MeCN/H2O 95:5 with 0.1 M NBu4PF6 0.1 M KHCO3/HCO−3 buffer pH 6.5 MeCN/H2O 95:5 with 0.1 M NBu4PF6 MeCN with 0.1 M NBu4PF6 MeCN with 0.1 M NBu4ClO4 DMF and 1.1 M H2O with 0.1 M NBu4PF6 DMF and 0.62 M PhOH with 0.1 M NBu4PF6

253 254 255 255 256 257 258 259

Fig. 23 Electroactive polymeric film generated from [Ru(bpy)(CO)2(Cl)2] under reducing conditions, which has the generic chemical structure [Ru0(bpy)(CO)2]n and contains metal−metal bonds. Adapted with permission from Chardon-Noblat, S.; Deronzier, A.; Ziessel, R.; Zsoldos, D. Selective Synthesis and Electrochemical Behavior of Trans(Cl)- and Cis(Cl)-[Ru(Bpy)(CO)2Cl2] Complexes (Bpy ¼ 2,2‘-Bipyridine). Comparative Studies of their Electrocatalytic Activity toward the Reduction of Carbon Dioxide. Inorg. Chem. 1997, 36, 5384–5389. Copyright 1997 American Chemical Society.

Fig. 24 Molecular structure of [Ru(bpy)(tpy)(CO)]2+ with atom labeling. Hydrogen atoms are omitted for clarity. Adapted from Nagao, H.; Mizukawa, T.; Tanaka, K. Carbon-Carbon Bond Formation in the Electrochemical Reduction of Carbon Dioxide Catalyzed by a Ruthenium Complex. Inorg. Chem. 1994, 33, 3415–3420. Copyright 1994 American Chemical Society.

Introduction to the Organometallic Chemistry of Carbon Dioxide

491

coupling between electrons in the reduced ligand framework and the Fe center was proposed to be key for the unprecedented activity, as were the changes in electronic structure that resulted from the strained square pyramidal environment enforced by the ligand framework. A major advance in the development of earth-abundant electrocatalysts was a report by Bourrez et al. in 2011, who reported the electrocatalytic activity of Mn(bpy)(CO)3Br for CO2 reduction to CO using H2O as a proton donor.253 Besides the inactivity under aprotic conditions, the major difference between Mn and Re was the propensity of Mn following one-electron reduction to form metal-metal dimers at rates which approached the diffusion limit.294 Kubiak and co-workers prevented the competing metal-metal dimerization reaction by using mesityl groups on the bpy ligand framework, enabling significant increases in activity at low overpotentials.295–297 Rochford and co-workers pushed this design principle further, adding pendent alkyl ethers as proton donor relays in a related bpy-based ligand derivative, resulting in the enhancement of catalysis at low overpotentials.298 In 2012, Kang et al. disclosed an Ir-based catalyst which showed 85% efficiency for formate production with water as a proton source in acetonitrile solution.254 This Ir-based catalyst was inspired by the work on CO2 hydrogenation discussed above in Fig. 17. Thermally, analogous Ir(III) trihydrides require a base to generate formate from CO2, but Kang et al. showed that electrochemically, 5% H2O/MeCN mixtures were suitable for electrocatalytically generating formate, although only dihydrides were invoked as plausible intermediates. Berben and co-workers followed shortly thereafter with an anionic Fe carbonyl cluster, [Fe4N(CO)12]−, which was eventually optimized to near quantitative efficiency for formate in both organic and aqueous solutions.255,299,300 In these clusters, the delocalization afforded by the metal-metal bonds were proposed to be key for achieving low overpotential reactions by diffusing added charge and minimizing reorganization penalties during key reaction steps. Inspired by the activity and selectivity of the Mn and Re catalysts, Kubiak and co-workers explored the electrocatalytic activity of M(bpy)(CO)4 (M ¼ Mo and W) complexes in 2014.256 Reduction of these M(0) complexes by two electrons accessed quantitative selectivity for CO2 reduction to CO under aprotic conditions in acetonitrile. CO loss from the tetracarbonyl did not occur until the second reduction; the tetracarbonyl monoanion was stable enough to be crystallized for characterization by X-ray diffraction methods. Overall, these results suggest that for the Mo and W to be active at lower overpotentials, M(0) states should be avoided because of the strong back-bonding possible with the potential CO product. In spite of the known benefits of arene complexes in thermal CO2 hydrogenation,301 reports of molecular electrocatalysts containing metal-arene species are relatively rare. Rosas-Hernández et al. disclosed a mononuclear Fe carbonyl species with a redox-active cyclopentadione-based ligand in 2017.257 This Fe complex produces CO with 96% Faradaic efficiency under aprotic conditions, with residual water or acetonitrile proposed to facilitate the production of water as a co-product. The same year, Roy et al. reported a series of cobalt cyclopentadienyl complexes which reduced CO2 to formate in DMF-H2O mixtures with high efficiency.258 A key component of the reaction mechanism was the participation of N bases in the secondary coordination sphere, which were proposed to facilitate several steps of the reaction. Metals early in the periodic table represent an ongoing challenge for molecular catalysts, although recent efforts suggest that proper ligand design can facilitate activity.302,303 In 2020 Hooe et al. reported the first known molecular Cr complex for CO2 reduction in DMF solution, obtaining CO as the exclusive product with phenol present as a weak acid.259 A subsequent computational study revealed that this selectivity was achieved through a Cr(II)(bpy  − ) active state, which resulted in a Cr center with decreased s-symmetric base character, enabling kinetic selectivity for CO2 over protonation to generate a Cr hydride.304 Consistent with this interpretation, earlier work by Nichols et al. on the comparable Fe-based system had shown formate could be produced with 85  2% efficiency in DMF with phenol as a proton donor, indicating that d electron count impacts nucleophilic character in this ligand set: in the catalytically active state Cr has more p-basic character and a preference for CO2 binding, whereas Fe has relatively more s-basic character and activates proton donors initially, rather than CO2.305,306

1.15.4.3

Carboxylation

Several examples of catalytic carboxylation reactions using CO2 were introduced in Section 1.15.3.2, however, there are additional reaction systems worth emphasizing here.16,35,208 It should be mentioned that the current industrial production of carboxylic acids relies primarily on carbonylation reactions, either in the presence of water or with subsequent oxidation steps, and the non-specific direct oxidation of hydrocarbons.307 Electrocatalytic carboxylation reactions using Ni complexes with CO2 as a substrate have been demonstrated previously for alkynes generating alkenyl Mg2+ carboxylate salts as the isolated product (Fig. 25).308–310 These reactions require a sacrificial Mg metal anode to concomitantly generate Mg(II) ions as the Ni(II) complex is reduced to Ni(0) at the cathode. The Mg(II) ions are proposed to complete the catalytic cycle by favoring dissociation of the carboxylate product from the Ni center by assisting NidO bond cleavage.308,311 Related carboxylation reactions for a more diverse substrate set have been demonstrated with Ni complexes using heterogeneous reducing agents or organometallic hydrocarbon substrates.104,312–316 More recently, Ackerman and co-workers have developed Co-based electrosynthetic systems with an extended substrate scope.317 Thermally, in addition to the alkylzinc transmetallation examples above, slurries with Zn and Mn metal can drive comparable carboxylation reactions using a variety of transition metal centers as catalysts.28,103,204–213

1.15.4.4

CO2 copolymerization

Inoue et al. reported the first example of the copolymerization of CO2 and epoxide substrates in 1969, using diethylzinc as a catalyst.318–320 Since this initial discovery, the class of polycarbonate polymers produced in this reaction have attracted significant

492

Introduction to the Organometallic Chemistry of Carbon Dioxide

Fig. 25 Proposed catalytic mechanism for electrochemical reductive carboxylation of alkenes to alkenyl carboxylate salts mediated by [Ni(bpy)3]2+. A sacrificial Mg anode supplies stoichiometric amounts of Mg2+ during every catalytic cycle.308

Fig. 26 Overview of copolymerization of propylene oxide or cyclohexene oxide with CO2 to produce polycarbone and/or ether linkages, as well as the cyclic carbonate byproduct. Adapted from Kozak, C. M.; Ambrose, K.; Anderson, T. S. Copolymerization of Carbon Dioxide and Epoxides by Metal Coordination Complexes. Coord. Chem. Rev. 2018, 376, 565–587.

attention, given the aforementioned interest in developing reactions where CO2 serves as a C1 feedstock (Fig. 26).25 Depending on the epoxide used, polymeric materials suitable for use as engineering thermoplastics and resins, packaging materials, interliners for safety glass, and in pyrotechnics can be produced.215,216 Inoue went on to publish the first discrete catalyst for this reaction, demonstrating that aluminum tetraphenylporphyrins could copolymerize CO2 and epoxides (ethylene oxide, propylene oxide, cyclohexene oxide).321 At 0.25% loading of the Al porphyrin catalyst w.r.t. propylene oxide at 8.1 bar CO2 pressure, polymers with a PDI ¼ 1.08 were obtained after 168 h. For these molecular catalysts, quaternary ammonium salts or triphenylphosphine were required additives. In 1995, Kruber and Dellar reported the first Cr-based catalyst for CO2 copolymerization, using a tetraarylporphyrin ligand.322 In this study, a variety of epoxides and additives were tested; N-donor ligands like N-methylimidazole or 4-dimethylaminopyridine were required co-catalytic additives for polymerization to occur. That same year, Darensbourg and Holtcamp reported the catalytic activity of Zn(II) phenoxides in this reaction, observing copolymerization of propylene oxide and CO2 at temperatures of 80  C and CO2 pressures of 54 atm over the course of ca. 70 h.323 Subsequent optimization through mechanistic studies of a series of Zn(II) phenoxides reported by Darensbourg et al. in 1999 resulted in significant improvements in molecular weight and polydispersity.324 The prior year (1998), Coates and co-workers reported that the alternating copolymerization of CO2 and epoxides could also be achieved using a b-diminate zinc catalyst, which could also produce polymers with narrow molecular weight distributions.325 These last catalysts were particularly active: the catalysts functioned at ambient temperatures with CO2 pressures as low as 6.8 atm. Following these advancements, Inoue and co-workers reported that Mn-based tetraarylporphyrin complexes could catalyze the copolymerization of CO2 and cyclohexene oxide at only 1 atm of pressure and at temperatures of 80  C without requiring a co-catalyst.326 This report also disclosed that Mn(salophen)(OAc) (salophen ¼ N,N0 -o-phenylenediaminebis(salicylideneiminato) dianion; OAc ¼ acetate) was inactive for polymerization at an elevated pressure of 50 atm. The first report of a [Cr(salen)Cl] (salen ¼ N,N0 -ethylenebis(salicyliminato) dianion) complex catalyzing a copolymerization reaction involving CO2 appeared in a patent from Jacobsen and co-workers in 2000.327 The following year, Nguyen and Paddock reported a series of active [Cr(salen)Cl] derivatives as catalysts for producing molecular carbonates from a series of epoxides and CO2 when co-catalytic amounts of 4-dimethylaminopyridine were added to the reaction mixture.328 Equilibrium displacement reactions of the chloride ligand bound to the starting species were proposed to facilitate the coordination of CO2, with a bimolecular interaction between CO2 and epoxide adducts facilitating the ring-expansion reaction (Fig. 27). In 2002, Darensbourg and Yarbrough reported the first example of the corresponding copolymerization using a Cr(salen)-based derivative.329 Using in situ NMR studies under pressure, the authors

Introduction to the Organometallic Chemistry of Carbon Dioxide

493

Fig. 27 Proposed catalytic cycle for the production of propylene carbonate from CO2 and propylene oxide. R ¼ Me, MeCl, CHCH2, hexyl, benzyl, phenyl. Adapted with permission from Paddock, R.L.; Nguyen, S.T. Chemical CO2 Fixation: Cr(III) Salen Complexes As Highly Efficient Catalysts for the Coupling of CO2 and Epoxides. J. Am. Chem. Soc. 2001, 123, 11498–11499. Copyright 2001 American Chemical Society.

established that the carbonate chain growth had a first-order concentration dependence on catalyst, suggesting that the bimolecular step reported by Nguyen and Paddock (expected to be second-order in [catalyst]) was involved only as an initiating step. In 2003, Coates and co-workers reported that analogous Co complexes were also active330 and Darensbourg and Billodeaux disclosed the first examples of the [Al(salen)]+ derivatives in 2005.331 Excluding limited reports of heterobimetallic systems where catalysis occurred at metal centers other than iron,332,333 it was not until 2011 that a discrete bimetallic iron-based catalyst was reported to copolymerize CO2 with epoxides.334 Under optimized conditions, Buchard et al. reported that the diiron species mediated 90% conversion of cyclohexene oxide and 98% of propylene oxide into the corresponding polycarbonate polymers at 1 atm of CO2 pressure and at temperatures of 80  C. In 2013, Nakano et al. reported the first example of a mononuclear Fe-based catalyst, using corrole derivatives to support the active site.335

1.15.5

Conclusions and perspective

In thermal CO2 functionalization, strategies to overcome the stability of the linear OCO p manifold are required. Disruption of the linear configuration can be achieved with strong nucleophiles by attack at the C atom. However, while low valent metal centers and metal amide complexes have the intrinsic activity required for this reaction pathway, their significant nucleophilicity can be indiscriminate, resulting in poor selectivity. The key challenge for direct catalytic CO2 activation is to generate active species which do not create strong bases with significant s character, given the resultant activity overlap with alternative substrates like H+. In electrocatalytic CO2 reduction, this may be achieved by using redox-active ligand frameworks, which delocalize added electron density, allowing the complex to act in a manner reminiscent of a p base, rather than the metal center acting as a localized s base. Alternatively, kinetic selectivity can be achieved by tuning reaction conditions with sufficient mechanistic understanding, as has been achieved in copolymerization reactions of epoxides and CO2. Although competing polyether chains or cyclic carbonate products are possible, the use of co-catalytic bases or bimetallic catalyst platforms and variable CO2 pressures can facilitate chain growth resulting in polymers with high monomer incorporation and low PDI values. Thermal hydrogenation strategies to generate formate salts from CO2 have successfully circumvented the thermodynamic penalty of producing formic acid directly at the expense of using stoichiometric amounts of added base. An appropriate mechanistic understanding of relevant hydricities can identify alternative solvent conditions where milder bases (B) can be used to drive the reaction (CO2 + H2 + B Ð HCO−2 + BH+) at high rates.336 The favorable thermodynamics of methanol production from CO2 hydrogenation suggest that significant advances in deployable technologies are feasible if decreased renewable H2 costs make them financially competitive. An area of significant future opportunity is the development of catalytic reactions for isocyanate production from CO2. The use of a metal center to instead catalytically activate primary amines for this reaction could enable the synthesis of carbamate precursors to important isocyanate substrates directly from CO2.337 Additionally, the currently energy-intensive production of urea shares conceptual overlap with isocyanate production from CO2 and would also benefit greatly from catalyst development.338 The broad summary here is meant to highlight the foundational principles and key advances in organometallic chemistry related to CO2. The isolation and structural verification of key CO2 adducts spurred an interest in chemical transformations, first stoichiometric and then catalytic. Given societal concerns about the implications of growing atmospheric concentrations of

494

Introduction to the Organometallic Chemistry of Carbon Dioxide

CO2 and the environmental burden of increasing anthropogenic emissions, the next decade is likely to see significant strides made in deployable CO2 conversion processes. From a coordination chemistry perspective, an understanding of how C]O bonds are polarized and broken will be required to develop better capture and conversion processes.

Acknowledgment The University of Virginia for generous financial and infrastructural support.

References 1. Lindsey, A. S.; Jeskey, H. The Kolbe-Schmitt Reaction. Chem. Rev. 1957, 57, 583–620. 2. Shakhashiri, B. Z. Burning of Magnesium. In Chemical Demonstrations, The University of Wisconsin Press, 1983; vol. 1; pp 38–39. 3. Li, Z.; Mayer, R. J.; Ofial, A. R.; Mayr, H. From Carbodiimides to Carbon Dioxide: Quantification of the Electrophilic Reactivities of Heteroallenes. J. Am. Chem. Soc. 2020, 142, 8383–8402. 4. Eisenberg, R.; Hendriksen, D. E. The Binding and Activation of Carbon Monoxide, Carbon Dioxide, and Nitric Oxide and Their Homogeneously Catalyzed Reactions. In Advances in Catalysis; Eley, D. D., Pines, H., Weez, P. B., Eds.; Academic Press, 1979; vol. 28; pp 79–172. 5. Cokoja, M.; Bruckmeier, C.; Rieger, B.; Herrmann, W. A.; Kühn, F. E. Transformation of Carbon Dioxide With Homogeneous Transition-Metal Catalysts: A Molecular Solution to a Global Challenge?Angew. Chem. Int. Ed. 2011, 50, 8510–8537. 6. Braunstein, P.; Matt, D.; Nobel, D. Reactions of Carbon Dioxide With Carbon-Carbon Bond Formation Catalyzed by Transition-Metal Complexes. Chem. Rev. 1988, 88, 747–764. 7. Behr, A. Carbon Dioxide as an Alternative C1 Synthetic Unit: Activation by Transition-Metal Complexes. Angew. Chem. Int. Ed. Engl. 1988, 27, 661–678. 8. Aresta, M.; Dibenedetto, A. Utilisation of CO2 as a Chemical Feedstock: Opportunities and Challenges. Dalton Trans. 2007, ;2975–2992. 9. Sakakura, T.; Choi, J.-C.; Yasuda, H. Transformation of Carbon Dioxide. Chem. Rev. 2007, 107, 2365–2387. 10. Benson, E. E.; Kubiak, C. P.; Sathrum, A. J.; Smieja, J. M. Electrocatalytic and Homogeneous Approaches to Conversion of CO2 to Liquid Fuels. Chem. Soc. Rev. 2009, 38, 89–99. 11. Kou, Y.; Nabetani, Y.; Masui, D.; Shimada, T.; Takagi, S.; Tachibana, H.; Inoue, H. Direct Detection of Key Reaction Intermediates in Photochemical CO2 Reduction Sensitized by a Rhenium Bipyridine Complex. J. Am. Chem. Soc. 2014, 136, 6021–6030. 12. Kondratenko, E. V.; Mul, G.; Baltrusaitis, J.; Larrazábal, G. O.; Pérez-Ramírez, J. Status and Perspectives of CO2 Conversion Into Fuels and Chemicals by Catalytic, Photocatalytic and Electrocatalytic Processes. Energ. Environ. Sci. 2013, 6, 3112–3135. 13. Grignard, B.; Gennen, S.; Jérôme, C.; Kleij, A. W.; Detrembleur, C. Advances in the Use of CO2 as a Renewable Feedstock for the Synthesis of Polymers. Chem. Soc. Rev. 2019, 48, 4466–4514. 14. Kleij, A. W.; North, M.; Urakawa, A. CO2 Catalysis. ChemSusChem 2017, 10, 1036–1038. 15. Tappe, N. A.; Reich, R. M.; D’Elia, V.; Kühn, F. E. Current Advances in the Catalytic Conversion of Carbon Dioxide by Molecular Catalysts: An Update. Dalton Trans. 2018, 47, 13281–13313. 16. Janis, L. Transition Metal Catalyzed Reactions of Carbon Dioxide and Other Heterocumulenes. Curr. Org. Chem. 2005, 9, 605–623. 17. Appel, A. M.; Bercaw, J. E.; Bocarsly, A. B.; Dobbek, H.; DuBois, D. L.; Dupuis, M.; Ferry, J. G.; Fujita, E.; Hille, R.; Kenis, P. J. A.; Kerfeld, C. A.; Morris, R. H.; Peden, C. H. F.; Portis, A. R.; Ragsdale, S. W.; Rauchfuss, T. B.; Reek, J. N. H.; Seefeldt, L. C.; Thauer, R. K.; Waldrop, G. L. Frontiers, Opportunities, and Challenges in Biochemical and Chemical Catalysis of CO2 Fixation. Chem. Rev. 2013, 113, 6621–6658. 18. Aresta, M.; Dibenedetto, A.; Angelini, A. Catalysis for the Valorization of Exhaust Carbon: From CO2 to Chemicals, Materials, and Fuels. Technological Use of CO2. Chem. Rev. 2014, 114, 1709–1742. 19. Nocera, D. G. Solar Fuels and Solar Chemicals Industry. Acc. Chem. Res. 2017, 50, 616–619. 20. Morris, A. J.; Meyer, G. J.; Fujita, E. Molecular Approaches to the Photocatalytic Reduction of Carbon Dioxide for Solar Fuels. Acc. Chem. Res. 2009, 42, 1983–1994. 21. Senftle, T. P.; Carter, E. A. The Holy Grail: Chemistry Enabling an Economically Viable CO2 Capture, Utilization, and Storage Strategy. Acc. Chem. Res. 2017, 50, 472–475. 22. Álvarez, A.; Bansode, A.; Urakawa, A.; Bavykina, A. V.; Wezendonk, T. A.; Makkee, M.; Gascon, J.; Kapteijn, F. Challenges in the Greener Production of Formates/Formic Acid, Methanol, and DME by Heterogeneously Catalyzed CO2 Hydrogenation Processes. Chem. Rev. 2017, 117, 9804–9838. 23. Zhao, G.; Huang, X.; Wang, X.; Wang, X. Progress in Catalyst Exploration for Heterogeneous CO2 Reduction and Utilization: A Critical Review. J. Mater. Chem. A 2017, 5, 21625–21649. 24. Greenblatt, J. B.; Miller, D. J.; Ager, J. W.; Houle, F. A.; Sharp, I. D. The Technical and Energetic Challenges of Separating (Photo)Electrochemical Carbon Dioxide Reduction Products. Joule 2018, 2, 381–420. 25. Kozak, C. M.; Ambrose, K.; Anderson, T. S. Copolymerization of Carbon Dioxide and Epoxides by Metal Coordination Complexes. Coord. Chem. Rev. 2018, 376, 565–587. 26. Kuramochi, Y.; Ishitani, O.; Ishida, H. Reaction Mechanisms of Catalytic Photochemical CO2 Reduction Using Re(I) and Ru(II) Complexes. Coord. Chem. Rev. 2018, 373, 333–356. 27. Sordakis, K.; Tang, C.; Vogt, L. K.; Junge, H.; Dyson, P. J.; Beller, M.; Laurenczy, G. Homogeneous Catalysis for Sustainable Hydrogen Storage in Formic Acid and Alcohols. Chem. Rev. 2018, 118, 372–433. 28. Fujihara, T.; Tsuji, Y. Carboxylation Reactions Using Carbon Dioxide as the C1 Source Via Catalytically Generated Allyl Metal Intermediates. Front. Chem. 2019, 7. 29. De, S.; Dokania, A.; Ramirez, A.; Gascon, J. Advances in the Design of Heterogeneous Catalysts and Thermocatalytic Processes for CO2 Utilization. ACS Catal. 2020, 10, 14147–14185. 30. Steinlechner, C.; Roesel, A. F.; Oberem, E.; Päpcke, A.; Rockstroh, N.; Gloaguen, F.; Lochbrunner, S.; Ludwig, R.; Spannenberg, A.; Junge, H.; Francke, R.; Beller, M. Selective Earth-Abundant System for CO2 Reduction: Comparing Photo- and Electrocatalytic Processes. ACS Catal. 2019, ;2091–2100. 31. Nichols, A. W.; Machan, C. W. Secondary-Sphere Effects in Molecular Electrocatalytic CO2 Reduction. Front. Chem. 2019, 7https://doi.org/10.3389/fchem.2019.00397. 32. Francke, R.; Schille, B.; Roemelt, M. Homogeneously Catalyzed Electroreduction of Carbon Dioxide—Methods, Mechanisms, and Catalysts. Chem. Rev. 2018, 118, 4631–4701. 33. Artz, J.; Müller, T. E.; Thenert, K.; Kleinekorte, J.; Meys, R.; Sternberg, A.; Bardow, A.; Leitner, W. Sustainable Conversion of Carbon Dioxide: An Integrated Review of Catalysis and Life Cycle Assessment. Chem. Rev. 2018, 118, 434–504. 34. Darensbourg, D. J.; Kudaroski, R. A. The Activation of Carbon Dioxide by Metai Complexes. In Advances in Organometallic Chemistry; Stone, F. G. A., West, R., Eds.; Academic Press, 1983; vol. 22; pp 129–168. 35. Walther, D. Homogeneous-Catalytic Reactions of Carbon Dioxide With Unsatureated Substrates, Reversible CO2-Carriers and Transcarboxylation Reactions. Coord. Chem. Rev. 1987, 79, 135–174.

Introduction to the Organometallic Chemistry of Carbon Dioxide

495

36. Gibson, D. H. The Organometallic Chemistry of Carbon Dioxide. Chem. Rev. 1996, 96, 2063–2096. 37. Mondal, B.; Song, J.; Neese, F.; Ye, S. Bio-Inspired Mechanistic Insights Into CO2 Reduction. Curr. Opin. Chem. Biol. 2015, 25, 103–109. 38. Lü, J.-M.; Rosokha, S. V.; Neretin, I. S.; Kochi, J. K. Quinones as Electron Acceptors. X-Ray Structures, Spectral (EPR, UV −vis) Characteristics and Electron-Transfer Reactivities of Their Reduced Anion Radicals as Separated vs Contact Ion Pairs. J. Am. Chem. Soc. 2006, 128, 16708–16719. 39. Hendon, C. H.; Tiana, D.; Murray, A. T.; Carbery, D. R.; Walsh, A. Helical Frontier Orbitals of Conjugated Linear Molecules. Chem. Sci. 2013, 4, 4278–4284. 40. Klaus, S.; Lehenmeier, M. W.; Anderson, C. E.; Rieger, B. Recent Advances in CO2/Epoxide Copolymerization—New Strategies and Cooperative Mechanisms. Coord. Chem. Rev. 2011, 255, 1460–1479. 41. Ema, T.; Miyazaki, Y.; Koyama, S.; Yano, Y.; Sakai, T. A Bifunctional Catalyst for Carbon Dioxide Fixation: Cooperative Double Activation of Epoxides for the Synthesis of Cyclic Carbonates. Chem. Commun. 2012, 48, 4489–4491. 42. Yang, G.-W.; Zhang, Y.-Y.; Xie, R.; Wu, G.-P. Scalable Bifunctional Organoboron Catalysts for Copolymerization of CO2 and Epoxides With Unprecedented Efficiency. J. Am. Chem. Soc. 2020, 142, 12245–12255. 43. Kai, T.; Zhou, M.; Duan, Z.; Henkelman, G. A.; Bard, A. J. Detection of CO2 – in the Electrochemical Reduction of Carbon Dioxide in N,N-Dimethylformamide by Scanning Electrochemical Microscopy. J. Am. Chem. Soc. 2017, 139, 18552–18557. 44. Palmer, D. A.; Van Eldik, R. The Chemistry of Metal Carbonato and Carbon Dioxide Complexes. Chem. Rev. 1983, 83, 651–731. 45. Gibson, D. H. Carbon Dioxide Coordination Chemistry: Metal Complexes and Surface-Bound Species. What Relationships?Coord. Chem. Rev. 1999, 185-186, 335–355. 46. Yin, X.; Moss, J. R. Recent Developments in the Activation of Carbon Dioxide by Metal Complexes. Coord. Chem. Rev. 1999, 181, 27–59. 47. Murphy, L. J.; Robertson, K. N.; Kemp, R. A.; Tuononen, H. M.; Clyburne, J. A. C. Structurally Simple Complexes of CO2. Chem. Commun. 2015, 51, 3942–3956. 48. Leitner, W. The Coordination Chemistry of Carbon Dioxide and Its Relevance for Catalysis: A Critical Survey. Coord. Chem. Rev. 1996, 153, 257–284. 49. Mascetti, J. Carbon Dioxide Coordination Chemistry and Reactivity of Coordinated CO2. In Carbon Dioxide as Chemical Feedstock, Wiley, 2010; pp 55–88. 50. Miller, J. D. Reactions of Coordinated Carbon Dioxide. In Reactions of Coordinated Ligands; Braterman, P. S., Ed.; Springer US: Boston, MA, 1989; vol. 2; pp 1–52. 51. Creutz, C. Chapter 2—Carbon Dioxide Binding to Transition-Metal Centers. In Electrochemical and Electrocatalytic Reactions of Carbon Dioxide; Sullivan, B. P., Ed.; Elsevier: Amsterdam, 1993; pp 19–67. 52. Vol’pin, M. E.; Kolomnikov, I. S.; Lobeeva, T. S. A rhodium-carbon dioxide complex. Bull. Acad. Sci. USSR 1969, 18, 1945. 53. Jolly, P. W.; Jonas, K.; Krüger, C.; Tsay, Y. H. The Preparation, Reactions and Structure of Bis[Bis(Tricyclohexylphosphine)Nickel] Dinitrogen, {[(C6H11)3P]2Ni}2N2. J. Organomet. Chem. 1971, 33, 109–122. 54. Aresta, M.; Nobile, C. F.; Albano, V. G.; Forni, E.; Manassero, M. New Nickel–Carbon Dioxide Complex: Synthesis, Properties, and Crystallographic Characterization of (Carbon Dioxide)-Bis(Tricyclohexylphosphine)Nickel. J. Chem. Soc. Chem. Commun. 1975, ;636–637. 55. Dohring, A.; Jolly, P. W.; Kruger, C.; Romão, M. J. The Ni(0)–CO2 System: Structure and Reactions of [Ni(PCy3)2(Z-CO2)]. Z. Naturforsh. B 1985, 40, 484. 56. Ovenall, D.; Whiffen, D. Electron Spin Resonance and Structure of the CO-2 Radical Ion. Mol. Phys. 1961, 4, 135–144. 57. Hartman, K. O.; Hisatsune, I. C. Infrared Spectrum of Carbon Dioxide Anion Radical. J. Chem. Phys. 1966, 44, 1913–1918. 58. Morton, J. R. Electron Spin Resonance Spectra of Oriented Radicals. Chem. Rev. 1964, 64, 453–471. 59. Bristow, G. S.; Hitchcock, P. B.; Lappert, M. F. A Novel Carbon Dioxide Complex: Synthesis and Crystal Structure of [Nb(Z-C5H4Me)2(CH2SiMe3)(Z2-CO2)]. J. Chem. Soc. Chem. Commun. 1981, ;1145–1146. 60. Gambarotta, S.; Arena, F.; Floriani, C.; Zanazzi, P. F. Carbon Dioxide Fixation: Bifunctional Complexes Containing Acidic and Basic Sites Working as Reversible Carriers. J. Am. Chem. Soc. 1982, 104, 5082–5092. 61. Calabrese, J. C.; Herskovitz, T.; Kinney, J. B. Carbon Dioxide Coordination Chemistry. 5. The Preparation and Structure of the Rhodium Complex Rh(.eta.1-CO2)(Cl)(diars)2. J. Am. Chem. Soc. 1983, 105, 5914–5915. 62. Herskovitz, T.; Parshall, G. W. Carbon Dioxide Complexes of Rh, Ir, Ni, Pd, and Pt. USA Patent. 63. Herskovitz, T. Carbon Dioxide Coordination Chemistry. 3. Adducts of Carbon Dioxide With Iridium(I) Complexes. J. Am. Chem. Soc. 1977, 99, 2391–2392. 64. Alvarez, R.; Carmona, E.; Poveda, M. L.; Sanchez-Delgado, R. Carbon Dioxide Chemistry. The Synthesis and Properties of Trans-Bis(Carbon Dioxide)Tetrakis(Trimethylphosphine) Molybdenum (Trans-[Mo(CO2)2(PMe3)4]): The First Stable Bis(Carbon Dioxide) Adduct of a Transition Metal. J. Am. Chem. Soc. 1984, 106, 2731–2732. 65. Gambarotta, S.; Floriani, C.; Chiesi-Villa, A.; Guastini, C. Carbon Dioxide and Formaldehyde Coordination on Molybdenocene to Metal and Hydrogen Bonds of the C1 Molecule in the Solid State. J. Am. Chem. Soc. 1985, 107, 2985–2986. 66. Alt, H. G.; Schwind, K.-H.; Rausch, M. D. Fixierung und aktivierung von CO2 und CS2 durch Cp2M(PM3)2-komplexe (Cp ¼ Z5-cyclopentadienyl; M ¼ Ti, Zr). J. Organomet. Chem. 1987, 321, C9–C12. 67. Ishida, T.; Hayashi, T.; Mizobe, Y.; Hidai, M. Preparation and Properties of Molybdenum and Tungsten Dinitrogen Complexes. 38. Hydrido-Carbonato, Hydrido-Carbamato, and Carbon Dioxide Complexes of Tungsten Derived from the Carbonyl-Dinitrogen Complex Trans-[W(CO)(N2)(Ph2PCH2CH2PPh2)2]. Inorg. Chem. 1992, 31, 4481–4485. 68. Alvarez, R.; Carmona, E.; Gutierrez-Puebla, E.; Marín, J. M.; Monge, A.; Poveda, M. L. Synthesis and X-Ray Crystal Structure of [Mo(CO2)2(PMe3)3(CNPri)]: The First Structurally Characterized Bis(Carbon Dioxide) Adduct of a Transition Metal. J. Chem. Soc. Chem. Commun. 1984, ;1326–1327. 69. Chatt, J.; Kubota, M.; Leigh, G. J.; March, F. C.; Mason, R.; Yarrow, D. J. A Possible Carbon Dioxide Complex of Molybdenum and Its Rearrangement Product di-mCarbonato-Bis{carbonyltris(dimethylphenylphosphine)molybdenum}: X-ray Crystal Structure. J. Chem. Soc. Chem. Commun. 1974, ;1033–1034. 70. Alvarez, R.; Carmona, E.; Marin, J. M.; Poveda, M. L.; Gutierrez-Puebla, E.; Monge, A. Carbon Dioxide Chemistry. Synthesis, Properties, and Structural Characterization of Stable Bis(carbon dioxide) Adducts of Molybdenum. J. Am. Chem. Soc. 1986, 108, 2286–2294. 71. Carmona, E.; Munoz, M. A.; Perez, P. J.; Poveda, M. L. Rotational Isomerism in Bis(Carbon Dioxide) Complexes of Molybdenum Generated by Conrotatory Motion of the CO2 Ligands. Organometallics 1990, 9, 1337–1339. 72. Carmona, E.; Hughes, A. K.; Munoz, M. A.; O’Hare, D. M.; Perez, P. J.; Poveda, M. L. Rotational Isomerism and Fluxional Behavior of Bis(Carbon Dioxide) Adducts of Molybdenum. J. Am. Chem. Soc. 1991, 113, 9210–9218. 73. Castro-Rodriguez, I.; Nakai, H.; Zakharov, L. N.; Rheingold, A. L.; Meyer, K. A Linear, O-Coordinated Z1-CO2 Bound to Uranium. Science 2004, 305, 1757–1759. 74. Viasus, C. J.; Gabidullin, B.; Gambarotta, S. Linear End-On Coordination Modes of CO2. Angew. Chem. Int. Ed. 2019, 58, 14887–14890. 75. Vaska, L. Catalytic Activation of Carbon Dioxide by Metal Complexes. J. Mol. Catal. 1988, 47, 381–388. 76. Darensbourg, D. J.; Bauch, C. G.; Ovalles, C. Metal-Induced Transformations of Carbon Dioxide. In Catalytic Activation of Carbon Dioxide, ACS Symposium Series American Chemical Society, 1988; vol. 363; pp 26–41. 77. Brunet, J. J. Tetracarbonylhydridoferrates, MHFe(CO)4: Versatile Tools in Organic Synthesis and Catalysis. Chem. Rev. 1990, 90, 1041–1059. 78. Darensbourg, D. J.; Rokicki, A.; Darensbourg, M. Y. Facile Reduction of Carbon Dioxide by Anionic Group 6b Metal Hydrides. Chemistry Relevant to Catalysis of the Water-Gas Shift Reaction. J. Am. Chem. Soc. 1981, 103, 3223–3224. 79. Sullivan, B. P.; Meyer, T. J. Kinetics and Mechanism of Carbon Dioxide Insertion Into a Metal-Hydride Bond. A Large Solvent Effect and an Inverse Kinetic Isotope Effect. Organometallics 1986, 5, 1500–1502. 80. Hieber, W.; Leutert, F. Über Metallcarbonyle. XII. Die Basenreaktion Des Eisenpentacarbonyls Und Die Bildung Des Eisencarbonylwasserstoffs. Z. Anorg. Allg. Chem. 1932, 204, 145–164. 81. Feigl, F.; Krumholz, P. Über die Einwirkung von Alkalialkoholaten auf Eisenpentakarbonyl. Monatsh. Chem. 1932, 59, 314–327. 82. Feigl, F.; Krumholz, P. Über salze des eisencarbonylwasserstoffs. Z. Anorg. Allg. Chem. 1933, 215, 242–248. 83. Krumholz, P.; Stettiner, H. The Acid Properties of Iron Tetracarbonyl Hydride. J. Am. Chem. Soc. 1949, 71, 3035–3039.

•

496

Introduction to the Organometallic Chemistry of Carbon Dioxide

84. Weinberger, B.; Tanguy, G.; Des Abbayes, H. A Mild Phase Transfer Synthesis of the Ylid Adduct (CO)4FeCH2P(C6H5)3 From Iron Pentacarbonyl and Dichloromethane: Evidence for the Transient Generation of the Tetracarbonyl Ferrate Anion Fe(CO)42 −. J. Organomet. Chem. 1985, 280, C31–C33. 85. Tanguy, G.; Clement, J.-C.; des Abbayes, H. A Mild Phase Transfer Synthesis of the m-Methylenebis-(Tetracarbonyliron) Complex m-CH2Fe2(CO)8 From Iron Pentacarbonyl and Dibromomethane. J. Organomet. Chem. 1986, 314, C43–C45. 86. Machan, C. W.; Sampson, M. D.; Kubiak, C. P. A Molecular Ruthenium Electrocatalyst for the Reduction of Carbon Dioxide to CO and Formate. J. Am. Chem. Soc. 2015, 137, 8564–8571. 87. Grice, N.; Kao, S. C.; Pettit, R. Chemical Properties of Metallocarboxylic Acids of Transition Metals. J. Am. Chem. Soc. 1979, 101, 1627–1628. 88. Sweet, J. R.; Graham, W. A. G. Hydrido and Hydroxycarbonyl Compounds of the Carbonyl(.eta.-cyclopentadienyl)Nitrosylrhenium Group. Organometallics 1982, 1, 982–986. 89. Schneck, F.; Ahrens, J.; Finger, M.; Stückl, A. C.; Würtele, C.; Schwarzer, D.; Schneider, S. The Elusive Abnormal CO2 Insertion Enabled by Metal-Ligand Cooperative Photochemical Selectivity Inversion. Nat. Commun. 2018, 9, 1161. 90. Kohnle, J. F.; Slaugh, L. H.; Nakamaye, K. L. Novel Effect of Carbon Dioxide on Catalyst Properties. Dimerization of butadiene. J. Am. Chem. Soc. 1969, 91, 5904–5905. 91. Sasaki, Y.; Inoue, Y.; Hashimoto, H. Reaction of Carbon Dioxide With Butadiene Catalysed by Palladium Complexes. Synthesis of 2-Ethylidenehept-5-en-4-Olide. J. Chem. Soc. Chem. Commun. 1976, ;605–606. 92. Musco, A. Co-Oligomerization of Butadiene and Carbon Dioxide Catalysed by Tertiary Phosphine–Palladium Complexes. J. Chem. Soc. Perkin Trans. 1980, 1, 693–698. 93. Inoue, Y.; Sasaki, Y.; Hashimoto, H. Incorporation of CO2 in Butadiene Dimerization Catalyzed by Palladium Complexes. Formation of 2-Ethylidene-5-Hepten-4-Olide. Bull. Chem. Soc. Jpn. 1978, 51, 2375–2378. 94. Musco, A.; Perego, C.; Tartiari, V. Telomerization Reactions of Butadiene and CO2 Catalyzed by Phosphine Pd(0) Complexes: (E)-2-Ethylidenehept-6-en-5-Olide and Octadienyl Esters of 2-Ethylidenehepta-4,6-Dienoic Acid. Inorg. Chim. Acta 1978, 28, L147–L148. 95. Behr, A.; Juszak, K. D.; Keim, W. Synthese Von 2-Ethyliden-6-Hepten-5-Olid. In Synthese de L’Ethylidene-2 Heptene-6olide-5, 1983. 96. Behr, A.; Juszak, K.-D. Palladium-Catalyzed Reaction of Butadiene and Carbon Dioxide. J. Organomet. Chem. 1983, 255, 263–268. 97. Behr, A.; He, R.; Juszak, K.-D.; Krüger, C.; Tsay, Y.-H. Steuerungsmöglichkeiten bei der übergangsmetall-katalysierten Umsetzung von 1,3-Dienen mit Kohlendioxid. Chem. Ber. 1986, 119, 991–1015. 98. Braunstein, P.; Matt, D.; Nobel, D. Carbon Dioxide Activation and Catalytic Lactone Synthesis by Telomerization of Butadiene and Carbon Dioxide. J. Am. Chem. Soc. 1988, 110, 3207–3212. 99. Jolly, P. W. Z3-Allylpalladium Compounds. Angew. Chem. Int. Ed. Engl. 1985, 24, 283–295. 100. Trinh, H. Reaktionen von Kohlendioxid und Schwefeldioxid mit j3-Allyl-Nickel-, Palladium- und Platin-Komplexen. 101. Schenker, G. Modellkomplexe Fuer Zwischenstufen PD-Katalysierter Reaktionen des Butadiens und Isoprens; Universitaet Bochum: Bochum, 1984. 102. Ukai, K.; Aoki, M.; Takaya, J.; Iwasawa, N. Rhodium(I)-Catalyzed Carboxylation of Aryl- and Alkenylboronic Esters with CO2. J. Am. Chem. Soc. 2006, 128, 8706–8707. 103. Tortajada, A.; Juliá-Hernández, F.; Börjesson, M.; Moragas, T.; Martin, R. Transition-Metal-Catalyzed Carboxylation Reactions With Carbon Dioxide. Angew. Chem. Int. Ed. 2018, 57, 15948–15982. 104. Yeung, C. S.; Dong, V. M. Beyond Aresta’s Complex: Ni- and Pd-Catalyzed Organozinc Coupling With CO2. J. Am. Chem. Soc. 2008, 130, 7826–7827. 105. Phapale, V. B.; Cárdenas, D. J. Nickel-Catalysed Negishi Cross-Coupling Reactions: Scope and Mechanisms. Chem. Soc. Rev. 2009, 38, 1598–1607. 106. Ochiai, H.; Jang, M.; Hirano, K.; Yorimitsu, H.; Oshima, K. Nickel-Catalyzed Carboxylation of Organozinc Reagents With CO2. Org. Lett. 2008, 10, 2681–2683. 107. Kobayashi, K.; Kondo, Y. Transition-Metal-Free Carboxylation of Organozinc Reagents Using CO2 in DMF Solvent. Org. Lett. 2009, 11, 2035–2037. 108. Ohishi, T.; Nishiura, M.; Hou, Z. Carboxylation of Organoboronic Esters Catalyzed by N-Heterocyclic Carbene Copper(I) Complexes. Angew. Chem. Int. Ed. 2008, 47, 5792–5795. 109. Takaya, J.; Tadami, S.; Ukai, K.; Iwasawa, N. Copper(I)-Catalyzed Carboxylation of Aryl- and Alkenylboronic Esters. Org. Lett. 2008, 10, 2697–2700. 110. Chisholm, M. H.; Extine, M. Tris(dimethylaminato)tris(N,N-dimethylcarbamato)tungsten(VI). Product of the Remarkable Reaction Between Hexakis(dimethylaminato)Tungsten and Carbon Dioxide. J. Am. Chem. Soc. 1974, 96, 6214–6216. 111. Broere, D. L. J.; Mercado, B. Q.; Holland, P. L. Selective Conversion of CO2 Into Isocyanate by Low-Coordinate Iron Complexes. Angew. Chem. Int. Ed. 2018, 57, 6507–6511. 112. Dickie, D. A.; Gislason, K. B.; Kemp, R. A. Formation of Phosphino-Substituted Isocyanate by Reaction of CO2 with Group 2 Complexes Based on the (Me3Si)(I-Pr2P)NH Ligand. Inorg. Chem. 2012, 51, 1162–1169. 113. Yin, H.; Carroll, P. J.; Schelter, E. J. Reactions of a Cerium(III) Amide With Heteroallenes: Insertion, Silyl-Migration and de-Insertion. Chem. Commun. 2016, 52, 9813–9816. 114. Krummenacher, I.; Cummins, C. C. Carbon–Phosphorus Triple Bond Formation Through Multiple Bond Metathesis of an Anionic Niobium Phosphide With Carbon Dioxide. Polyhedron 2012, 32, 10–13. 115. Maria, L.; Bandeira, N. A. G.; Marçalo, J.; Santos, I. C.; Gibson, J. K. CO2 Conversion to Phenyl Isocyanates by Uranium(VI) Bis(Imido) Complexes. Chem. Commun. 2020, 56, 431–434. 116. Cleaves, P. A.; Kefalidis, C. E.; Gardner, B. M.; Tuna, F.; McInnes, E. J. L.; Lewis, W.; Maron, L.; Liddle, S. T. Terminal Uranium(V/VI) Nitride Activation of Carbon Dioxide and Carbon Disulfide: Factors Governing Diverse and Well-Defined Cleavage and Redox Reactions. Chem. A Eur. J. 2017, 23, 2950–2959. 117. Keane, A. J.; Farrell, W. S.; Yonke, B. L.; Zavalij, P. Y.; Sita, L. R. Metal-Mediated Production of Isocyanates, R3ENCO from Dinitrogen, Carbon Dioxide, and R3ECl. Angew. Chem. Int. Ed. 2015, 54, 10220–10224. 118. Sobota, P.; Jez˙owska-Trzebiatowska, B.; Janas, Z. Insertion of CO2 into the Metal—Nitrogen Bond Formed in the Reaction With Molecular Nitrogen. J. Organomet. Chem. 1976, 118, 253–258. 119. Chandra, G.; Jenkins, A. D.; Lappert, M. F.; Srivastava, R. C. Amido-Derivatives of Metals and Metalloids. Part X. Reactions of Titanium(IV), Zirconium(IV), and Hafnium(IV) Amides With Unsaturated Substrates, and Some Related Experiments With Amides of Boron, Silicon, Germanium, and Tin(IV). J. Chem. Soc. A 1970, ;2550–2558. 120. Tang, Y.; Zakharov, L. N.; Rheingold, A. L.; Kemp, R. A. Insertion of Carbon Dioxide Into Mg− N Bonds. Structural Characterization of a Previously Unknown Z2 Chelation Mode to Magnesium in Magnesium Carbamates. Organometallics 2004, 23, 4788–4791. 121. Chisholm, M. H.; Extine, M. Carbon Dioxide Exchange Reactions Involving Early Transition-Metal NN-Dimethylcarbamato Compounds: Reversible Insertion of Carbon Dioxide into Transition-Metal–Nitrogen s-Bonds. J. Chem. Soc. Chem. Commun. 1975, ;438–439. 122. Dalton, R. F.; Jones, K. Reactions of Metal–Nitrogen Compounds With Unsaturated Substrates. Part I. Reactions of Aminostannanes, Stannazanes, And Stannylamines With Carbon Dioxide, Carbon Disulphide, and Carbonyl Sulphide. J. Chem. Soc. A 1970, ;590–594. 123. Chisholm, M. H.; Extine, M. Reactions of Transition Metal-Nitrogen.Sigma.-Bonds. II. Pentakis(N,N-Dimethylcarbamato)Niobium(V) and Its Facile Exchange Reaction With Carbon Dioxide. J. Am. Chem. Soc. 1975, 97, 1623–1625. 124. Chisholm, M. H.; Extine, M. New Metalloorganic Compounds of Tungsten(III). J. Am. Chem. Soc. 1975, 97, 5625–5627. 125. Chisholm, M. H.; Extine, M. W. Reactions of Transition Metal-Nitrogen. sigma. bonds. 3. Early Transition Metal N,N-Dimethylcarbamates. Preparation, Properties, and Carbon Dioxide Exchange Reactions. J. Am. Chem. Soc. 1977, 99, 782–792. 126. Breederveld, H. The Reaction of Dialkylaminosilanes With Carbon Dioxide and With Carbon Disulphide. Recl. Trav. Chim. Pays-Bas 1962, 81, 276–278. 127. Oertel, G.; Malz, H.; Holtschmidt, H. Addition Reactions to Amides of Silicon, Phosphorus, Arsenic, and Sulfur. Chem. Ber. 1964, 97, 891–902. 128. Cragg, R. H.; Lappert, M. F. Amino-derivatives of metals and metalloids. Part IV. Aminosilylation and aminophosphination of some unsaturated substrates. J. Chem. Soc. A 1966, ;82–85. 129. Peterson, L.; Thé, K. The Reactions of Substituted Germyl-and Silylhydrazines With BX3 and CY2 Acceptor Species. Can. J. Chem. 1972, 50, 562–566. 130. Satge, J.; Lesbre, M.; Baudet, M. Transamination Reactions of Trialkylgermylamines. Compt. Rend. 1964, 259, 4733. 131. Rivière-Baudet, M.; Satgé, J. Réactions d’insertions sur la liaison germanium-azote de quelques trialcoylgermylamines. Bull. Soc. Chim. Fr. 1969, ;1356.

Introduction to the Organometallic Chemistry of Carbon Dioxide

497

132. George, T. A.; Jones, K.; Lappert, M. F. 385. Amino-Derivatives of Metals and Metalloids. Part II. Aminostannylation of Unsaturated Substrates, and the Infrared Spectra and Structures of Carbamato- and Dithiocarbamato-Trimethylstannanes and Related Compounds. J. Chem. Soc. 1965, ;2157–2165. 133. Bloodworth, A. J.; Davies, A. G.; Vasishtha, S. C. Organometallic Reactions. Part VII. Further addition Reactions of Tributyltin Methoxide and of Bistributyltin Oxide. J. Chem. Soc. C 1967, ;1309–1313. 134. Davies, A. G.; Harrison, P. G. Organometallic Reactions. Part VIII. Addition Reactions of Dibutyltin Dimethoxide and Related Compounds. J. Chem. Soc. C 1967, ;1313–1317. 135. Koketsu, J.; Ishii, Y. Reactions of Tris(Dimethylamino)Stibine With Organic Compounds. J. Chem. Soc. C 1971, ;511–513. 136. Noltes, J. G. Studies in Group IV Organometallic Chemistry. XIX: The Reaction of Trialkyltin Alkoxides With (Thio)Acylamides and (Thio)Carbamates. Recl. Trav. Chim. Pays-Bas 1965, 84, 799–805. 137. Inoue, S.; Yokoo, Y. Reaction of Organoaluminum Coordination Compound With Carbon Dioxide. Bull. Chem. Soc. Jpn. 1972, 45, 3651–3653. 138. Yoshida, Y.; Ishii, S.; Kawato, A.; Yamashita, T.; Yano, M.; Inoue, S. Novel Syntheses of Arylcarbamic Esters From Carbon Dioxide and Aromatic Amine Via a Zinc Carbamate. Bull. Chem. Soc. Jpn. 1988, 61, 2913–2916. 139. Dell’Amico, D. B.; Calderazzo, F.; Labella, L.; Marchetti, F.; Pampaloni, G. Converting Carbon Dioxide into Carbamato Derivatives. Chem. Rev. 2003, 103, 3857–3898. 140. Bresciani, G.; Biancalana, L.; Pampaloni, G.; Marchetti, F. Recent Advances in the Chemistry of Metal Carbamates. Molecules 2020, 25, 3603. 141. Gambhir, A.; Tavoni, M. Direct Air Carbon Capture and Sequestration: How it Works and How it Could Contribute to Climate-Change Mitigation. One Earth 2019, 1, 405–409. 142. Mahmoudkhani, M.; Keith, D. W. Low-Energy Sodium Hydroxide Recovery for CO2 Capture From Atmospheric Air—Thermodynamic Analysis. Int. J. Greenhouse Gas Control 2009, 3, 376–384. 143. Stolaroff, J. K.; Keith, D. W.; Lowry, G. V. Carbon Dioxide Capture From Atmospheric Air Using Sodium Hydroxide Spray. Environ. Sci. Technol. 2008, 42, 2728–2735. 144. Baciocchi, R.; Storti, G.; Mazzotti, M. Process Design and Energy Requirements for the Capture of Carbon Dioxide From Air. Chem. Eng. Process. Process Intensif. 2006, 45, 1047–1058. 145. Christianson, D. W.; Fierke, C. A. Carbonic Anhydrase: Evolution of the Zinc Binding Site by Nature and by Design. Acc. Chem. Res. 1996, 29, 331–339. 146. Frost, S. C.; McKenna, R. Carbonic Anhydrase: Mechanism, Regulation, Links to Disease, and Industrial Applications; 75. Springer Science & Business Media, 2013. 147. Supuran, C.; De Simone, G. Carbonic Anhydrases as Biocatalysts: From Theory to Medical and Industrial Applications; Elsevier, 2015. 148. Looney, A.; Parkin, G.; Alsfasser, R.; Ruf, M.; Vahrenkamp, H. Zinc Pyrazolylborate Complexes Relevant to the Biological Function of Carbonic Anhydrase. Angew. Chem. Int. Ed. Engl. 1992, 31, 92–93. 149. Looney, A.; Han, R.; McNeill, K.; Parkin, G. Tris(Pyrazolyl)Hydroboratozinc Hydroxide Complexes as Functional Models for Carbonic Anhydrase: On the Nature of the Bicarbonate Intermediate. J. Am. Chem. Soc. 1993, 115, 4690–4697. 150. Alsfasser, R.; Trofimenko, S.; Looney, A.; Parkin, G.; Vahrenkamp, H. A Mononuclear Zinc Hydroxide Complex Stabilized by a Highly Substituted Tris(Pyrazolyl)Hydroborato Ligand: Analogies With the Enzyme Carbonic Anhydrase. Inorg. Chem. 1991, 30, 4098–4100. 151. Parkin, G. The Bioinorganic Chemistry of Zinc: Synthetic Analogues of Zinc Enzymes That Feature Tripodal Ligands. Chem. Commun. 2000, ;1971–1985. 152. Trofimenko, S. Scorpionates; Imperial College Press, 1999. 153. Armstrong, W. H.; Lippard, S. J. (.mu.-Oxo)bis(.mu.-acetato)bis(tri-1-pyrazolylborato)diiron(III), [(HBpz3)FeO(CH3CO2)2Fe(HBpz3)]: Model for the Binuclear Iron Center of Hemerythrin. J. Am. Chem. Soc. 1983, 105, 4837–4838. 154. Burzlaff, N. Biomimetic Trispyrazolylborato Iron Complexes. Angew. Chem. Int. Ed. 2009, 48, 5580–5582. 155. Kimura, E.; Shiota, T.; Koike, T.; Shiro, M.; Kodama, M. A Zinc(II) Complex of 1,5,9-Triazacyclododecane ([12]aneN3) as a Model for Carbonic Anhydrase. J. Am. Chem. Soc. 1990, 112, 5805–5811. 156. Kimura, E. Macrocyclic Polyamines With Intelligent Functions. Tetrahedron 1992, 48, 6175–6217. 157. Kimura, E. Model Studies for Molecular Recognition of Carbonic Anhydrase and Carboxypeptidase. Acc. Chem. Res. 2001, 34, 171–179. 158. Zhang, X.; van Eldik, R.; Koike, T.; Kimura, E. Kinetics and Mechanism of the Hydration of Carbon Dioxide and Dehydration of Bicarbonate Catalyzed by a Zinc (II) Complex of 1,5,9-Triazacyclododecane as a Model for Carbonic Anhydrase. Inorg. Chem. 1993, 32, 5749–5755. 159. Floyd, W. C.; Baker, S. E.; Valdez, C. A.; Stolaroff, J. K.; Bearinger, J. P.; Satcher, J. H.; Aines, R. D. Evaluation of a Carbonic Anhydrase Mimic for Industrial Carbon Capture. Environ. Sci. Technol. 2013, 47, 10049–10055. 160. Morimoto, T.; Nakajima, T.; Sawa, S.; Nakanishi, R.; Imori, D.; Ishitani, O. CO2 Capture by a Rhenium(I) Complex With the Aid of Triethanolamine. J. Am. Chem. Soc. 2013, 135, 16825–16828. 161. Koizumi, H.; Chiba, H.; Sugihara, A.; Iwamura, M.; Nozaki, K.; Ishitani, O. CO2 Capture by Mn(I) and Re(I) Complexes With a Deprotonated Triethanolamine Ligand. Chem. Sci. 2019, 10, 3080–3088. 162. Kumagai, H.; Nishikawa, T.; Koizumi, H.; Yatsu, T.; Sahara, G.; Yamazaki, Y.; Tamaki, Y.; Ishitani, O. Electrocatalytic Reduction of Low Concentration CO2. Chem. Sci. 2019, 10, 1597–1606. 163. Bigorgne, M.; Rassat, L. Étude spectrographique des liaisons C–O et C–N des complexes-CNj des nickel et molybdène carbonyle. Bull. Soc. Chim. Fr. 1963, 5, 295–303. 164. Kim, W. Y.; Chang, J.-S.; Park, S.-E.; Ferrence, G.; Kubiak, C. P. Mechanistic and IR Spectroelectrochemical Studies for Alkali Metal Ion Catalyzed Multiple Bond Metathesis Reactions of Carbon Dioxide. Chem. Lett. 1998, 27, 1063–1064. 165. DeLaet, D. L.; Fanwick, P. E.; Kubiak, C. P. Carbon Dioxide Induced Metathesis of C^N and C^O Triple Bonds of Methyl Isocyanide and Carbon Monoxide. J. Chem. Soc. Chem. Commun. 1987, ;1412–1413. 166. DeLaet, D. L.; Del Rosario, R.; Fanwick, P. E.; Kubiak, C. P. Carbon Dioxide Chemistry and Electrochemistry of a Binuclear Cradle Complex of Nickel(0), Ni2(.mu.-CNMe)(CNMe)2 (PPh2CH2PPh2)2. J. Am. Chem. Soc. 1987, 109, 754–758. 167. Lemke, F. R.; DeLaet, D. L.; Gao, J.; Kubiak, C. P. Photochemical Activation of Carbon Dioxide. Transient Absorbance Kinetic Studies of the Addition of Carbon Dioxide to a Metal-to-Bridging Ligand Charge-Transfer State of a Binuclear Nickel(0) Complex. J. Am. Chem. Soc. 1988, 110, 6904–6906. 168. DeLaet, D. L.; Fanwick, P. E.; Kubiak, C. P. Binuclear Complexes of Nickel(0): Comparison of a Bridging Methyl Isocyanide and a Bridging (Methylamino)Carbyne Ligand. Organometallics 1986, 5, 1807–1811. 169. Sita, L. R.; Babcock, J. R.; Xi, R. Facile Metathetical Exchange Between Carbon Dioxide and the Divalent Group 14 Bisamides M[N(SiMe3)2]2 (M ¼ Ge and Sn). J. Am. Chem. Soc. 1996, 118, 10912–10913. 170. Rimo, X.; Sita, L. R. Mechanistic Details for Metathetical Exchange Between XCO (X ¼ O and RN) and the Tin(II) Dimer, {Sn[N(SiMe3)2] (m-OBut)}2. Inorg. Chim. Acta 1998, 270, 118–122. 171. Ghosh, R.; Samuelson, A. G. Catalytic Metathesis of Carbon Dioxide With Heterocumulenes Mediated by Titanium Isopropoxide. Chem. Commun. 2005, ;2017–2019. 172. Kilgore, U. J.; Basuli, F.; Huffman, J. C.; Mindiola, D. J. Aryl Isocyanate, Carbodiimide, and Isocyanide Prepared From Carbon Dioxide. A Metathetical Group-Transfer Tale Involving a Titanium− Imide Zwitterion. Inorg. Chem. 2006, 45, 487–489. 173. Mokhtarzadeh, C. C.; Moore, C. E.; Rheingold, A. L.; Figueroa, J. S. Terminal Iron Carbyne Complexes Derived From Arrested CO2 Reductive Disproportionation. Angew. Chem. Int. Ed. 2017, 56, 10894–10899. 174. Savéant, J.-M. Molecular Catalysis of Electrochemical Reactions. Mechanistic Aspects. Chem. Rev. 2008, 108, 2348–2378. 175. Hull, J. F.; Himeda, Y.; Wang, W. H.; Hashiguchi, B.; Periana, R.; Szalda, D. J.; Muckerman, J. T.; Fujita, E. Reversible Hydrogen Storage Using CO2 and a Proton-Switchable Iridium Catalyst in Aqueous Media Under Mild Temperatures and Pressures. Nat. Chem. 2012, 4, 383–388. 176. Grasemann, M.; Laurenczy, G. Formic Acid as a Hydrogen Source—Recent Developments and Future Trends. Energ. Environ. Sci. 2012, 5, 8171–8181. 177. Li, J. J. Eschweiler–Clarke Reductive Alkylation of Amines. In Name Reactions: A Collection of Detailed Mechanisms and Synthetic Applications, Springer Berlin Heidelberg: Berlin, Heidelberg, 2009; pp 210–211.

498 178. 179. 180. 181. 182. 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.

Introduction to the Organometallic Chemistry of Carbon Dioxide Gladiali, S.; Alberico, E. Asymmetric Transfer Hydrogenation: Chiral Ligands and Applications. Chem. Soc. Rev. 2006, 35, 226–236. Fischer, F.; Tropsch, H. Über die direkte Synthese von Erdöl-Kohlenwasserstoffen bei gewöhnlichem Druck. (Erste Mitteilung). Ber. Dtsch. Chem. Ges. 1926, 59, 830–831. Fischer, F.; Tropsch, H. Über die direkte Synthese von Erdöl-Kohlenwasserstoffen bei gewöhnlichem Druck. (Zweite Mitteilung). Ber. Dtsch. Chem. Ges. 1926, 59, 832–836. Vosloo, A. C. Fischer–Tropsch: A Futuristic View. Fuel Process. Technol. 2001, 71, 149–155. Maitlis, P. M.; Zanotti, V. The Role of Electrophilic Species in the Fischer–Tropsch Reaction. Chem. Commun. 2009, ;1619–1634. Ojeda, M.; Nabar, R.; Nilekar, A. U.; Ishikawa, A.; Mavrikakis, M.; Iglesia, E. CO Activation Pathways and the Mechanism of Fischer–Tropsch Synthesis. J. Catal. 2010, 272, 287–297. Zhang, Q.; Kang, J.; Wang, Y. Development of Novel Catalysts for Fischer–Tropsch Synthesis: Tuning the Product Selectivity. ChemCatChem 2010, 2, 1030–1058. West, N. M.; Miller, A. J. M.; Labinger, J. A.; Bercaw, J. E. Homogeneous Syngas Conversion. Coord. Chem. Rev. 2011, 255, 881–898. Jiao, F.; Li, J.; Pan, X.; Xiao, J.; Li, H.; Ma, H.; Wei, M.; Pan, Y.; Zhou, Z.; Li, M.; Miao, S.; Li, J.; Zhu, Y.; Xiao, D.; He, T.; Yang, J.; Qi, F.; Fu, Q.; Bao, X. Selective Conversion of Syngas to Light Olefins. Science 2016, 351, 1065–1068. Hood, D. M.; Johnson, R. A.; Carpenter, A. E.; Younker, J. M.; Vinyard, D. J.; Stanley, G. G. Highly Active Cationic Cobalt(II) Hydroformylation Catalysts. Science 2020, 367, 542–548. Delolo, F. G.; dos Santos, E. N.; Gusevskaya, E. V. Anisole: A Further Step to Sustainable Hydroformylation. Green Chem. 2019, 21, 1091–1098. Vorholt, A. J. Hydroformylation of Renewables. In Homogeneous Catalysis With Renewables; Gaide, T., Behr, A., Behr, A., Vorholt, A. J., Eds.; Springer International Publishing: Cham, 2017; pp 41–64. Kalck, P.; Le Berre, C.; Serp, P. Recent Advances in the Methanol Carbonylation Reaction Into Acetic Acid. Coord. Chem. Rev. 2020, 402, 213078. Jessop, P. G.; Joó, F.; Tai, C.-C. Recent Advances in the Homogeneous Hydrogenation of Carbon Dioxide. Coord. Chem. Rev. 2004, 248, 2425–2442. Wang, W.-H.; Himeda, Y.; Muckerman, J. T.; Manbeck, G. F.; Fujita, E. CO2 Hydrogenation to Formate and Methanol as an Alternative to Photo- and Electrochemical CO2 Reduction. Chem. Rev. 2015, 115, 12936–12973. Garza, A. J.; Bell, A. T.; Head-Gordon, M. Mechanism of CO2 Reduction at Copper Surfaces: Pathways to C2 Products. ACS Catal. 2018, 8, 1490–1499. Zheng, Y.; Vasileff, A.; Zhou, X.; Jiao, Y.; Jaroniec, M.; Qiao, S.-Z. Understanding the Roadmap for Electrochemical Reduction of CO2 to Multi-Carbon Oxygenates and Hydrocarbons on Copper-Based Catalysts. J. Am. Chem. Soc. 2019, 141, 7646–7659. Yoshio, H.; Katsuhei, K.; Shin, S. Production of CO and CH4 in Electrochemical Reduction of CO2 at Metal Electrodes in Aqueous Hydrogen Carbonate Solution. Chem. Lett. 1985, 14, 1695–1698. Hori, Y.; Kikuchi, K.; Murata, A.; Suzuki, S. Production of Methane and Ethylene in Electrochemical Reduction of Carbon Dioxide at Copper Electrode in Aqueous Hydrogen Carbonate Solution. Chem. Lett. 1986, 15, 897–898. Angamuthu, R.; Byers, P.; Lutz, M.; Spek, A. L.; Bouwman, E. Electrocatalytic CO2 Conversion to Oxalate by a Copper Complex. Science 2010, 327, 313–315. Pokharel, U. R.; Fronczek, F. R.; Maverick, A. W. Reduction of Carbon Dioxide to Oxalate by a Binuclear Copper Complex. Nat. Commun. 2014, 5, 5883. Huber, G. W.; Iborra, S.; Corma, A. Synthesis of Transportation Fuels From Biomass: Chemistry, Catalysts, and Engineering. Chem. Rev. 2006, 106, 4044–4098. Lahijani, P.; Zainal, Z. A.; Mohammadi, M.; Mohamed, A. R. Conversion of the Greenhouse Gas CO2 to the Fuel Gas CO Via the Boudouard Reaction: A Review. Renew. Sustain. Energy Rev. 2015, 41, 615–632. Liu, A.; Gao, M.; Ren, X.; Meng, F.; Yang, Y.; Gao, L.; Yang, Q.; Ma, T. Current Progress in Electrocatalytic Carbon Dioxide Reduction to Fuels on Heterogeneous Catalysts. J. Mater. Chem. A 2020, 8, 3541–3562. Na, J.; Seo, B.; Kim, J.; Lee, C. W.; Lee, H.; Hwang, Y. J.; Min, B. K.; Lee, D. K.; Oh, H.-S.; Lee, U. General Technoeconomic Analysis for Electrochemical Coproduction Coupling Carbon Dioxide Reduction With Organic Oxidation. Nat. Commun. 2019, 10, 5193. Aresta, M., Ed.; In Carbon Dioxide as Chemical Feedstock; John Wiley & Sons, 2010. Huang, K.; Sun, C.-L.; Shi, Z.-J. Transition-Metal-Catalyzed C–C Bond Formation Through the Fixation of Carbon Dioxide. Chem. Soc. Rev. 2011, 40, 2435–2452. Tsuji, Y.; Fujihara, T. Carbon Dioxide as a Carbon Source in Organic Transformation: Carbon–Carbon Bond Forming Reactions by Transition-Metal Catalysts. Chem. Commun. 2012, 48, 9956–9964. Cai, X.; Xie, B. Direct Carboxylative Reactions for the Transformation of Carbon Dioxide Into Carboxylic Acids and Derivatives. Synthesis 2013, 45, 3305–3324. Liu, Q.; Wu, L.; Jackstell, R.; Beller, M. Using Carbon Dioxide as a Building Block in Organic Synthesis. Nat. Commun. 2015, 6, 5933. Yu, D.; Teong, S. P.; Zhang, Y. Transition Metal Complex Catalyzed Carboxylation Reactions with CO2. Coord. Chem. Rev. 2015, 293-294, 279–291. Börjesson, M.; Moragas, T.; Gallego, D.; Martin, R. Metal-Catalyzed Carboxylation of Organic (Pseudo)Halides With CO2. ACS Catal. 2016, 6, 6739–6749. Sekine, K.; Yamada, T. Silver-Catalyzed Carboxylation. Chem. Soc. Rev. 2016, 45, 4524–4532. Wang, S.; Du, G.; Xi, C. Copper-Catalyzed Carboxylation Reactions Using Carbon Dioxide. Org. Biomol. Chem. 2016, 14, 3666–3676. Chen, Y.-G.; Xu, X.-T.; Zhang, K.; Li, Y.-Q.; Zhang, L.-P.; Fang, P.; Mei, T.-S. Transition-Metal-Catalyzed Carboxylation of Organic Halides and Their Surrogates With Carbon Dioxide. Synthesis 2018, 50, 35–48. Luan, Y.-X.; Ye, M. Transition Metal-Mediated or Catalyzed Hydrocarboxylation of Olefins With CO2. Tetrahedron Lett. 2018, 59, 853–861. Darensbourg, D. J.; Holtcamp, M. W. Catalysts for the Reactions of Epoxides and Carbon Dioxide. Coord. Chem. Rev. 1996, 153, 155–174. Coates, G. W.; Moore, D. R. Discrete Metal-Based Catalysts for the Copolymerization of CO2 and Epoxides: Discovery, Reactivity, Optimization, and Mechanism. Angew. Chem. Int. Ed. 2004, 43, 6618–6639. Darensbourg, D. J. Making Plastics From Carbon Dioxide: Salen Metal Complexes as Catalysts for the Production of Polycarbonates From Epoxides and CO2. Chem. Rev. 2007, 107, 2388–2410. Sugimoto, H.; Inoue, S. Copolymerization of Carbon Dioxide and Epoxide. J. Polym. Sci. A Polym. Chem. 2004, 42, 5561–5573. Ra, E. C.; Kim, K. Y.; Kim, E. H.; Lee, H.; An, K.; Lee, J. S. Recycling Carbon Dioxide Through Catalytic Hydrogenation: Recent Key Developments and Perspectives. ACS Catal. 2020, 10, 11318–11345. Inoue, Y.; Izumida, H.; Sasaki, Y.; Hashimoto, H. Catalytic Fixation of Carbon Dioxide to Formic Acid by Transition-Metal Complexes under Mild Conditions. Chem. Lett. 1976, 5, 863–864. Jessop, P. G.; Ikariya, T.; Noyori, R. Homogeneous Hydrogenation of Carbon Dioxide. Chem. Rev. 1995, 95, 259–272. Kudo, K.; Phala, H.; Sugita, N.; Takezaki, Y. Synthesis of Dimethyl FORMAMIDE From Carbon Dioxide, Hydrogen and Dimethyl Amine Catalyzed by Palladium(II) Chloride. Chem. Lett. 1977, 6, 1495–1496. Kudo, K.; Sugita, N.; Takezaki, Y. Kinetic Study on the Synthesis of Alkali Formate From Carbon Dioxide and Hydrogen Catalyzed by Palladium(II) Chloride in an Aqueous Alkali Solution. Nippon Kagaku Kaishi 1977, 1977, 302–309. Leitner, W. Carbon Dioxide as a Raw Material: The Synthesis of Formic Acid and Its Derivatives from CO2. Angew. Chem. Int. Ed. Engl. 1995, 34, 2207–2221. Wesselbaum, S.; Hintermair, U.; Leitner, W. Continuous-Flow Hydrogenation of Carbon Dioxide to Pure Formic Acid Using an Integrated scCO2 Process With Immobilized Catalyst and Base. Angew. Chem. Int. Ed. 2012, 51, 8585–8588. Gassner, F.; Leitner, W. Hydrogenation of Carbon Dioxide to Formic Acid Using Water-Soluble Rhodium Catalysts. J. Chem. Soc. Chem. Commun. 1993, ;1465–1466. Jessop, P. G.; Ikariya, T.; Noyori, R. Homogeneous Catalytic Hydrogenation of Supercritical Carbon Dioxide. Nature 1994, 368, 231–233. Jessop, P. G.; Hsiao, Y.; Ikariya, T.; Noyori, R. Homogeneous Catalysis in Supercritical Fluids: Hydrogenation of Supercritical Carbon Dioxide to Formic Acid, Alkyl Formates, and Formamides. J. Am. Chem. Soc. 1996, 118, 344–355. Tanaka, R.; Yamashita, M.; Nozaki, K. Catalytic Hydrogenation of Carbon Dioxide Using Ir(III)−Pincer Complexes. J. Am. Chem. Soc. 2009, 131, 14168–14169.

Introduction to the Organometallic Chemistry of Carbon Dioxide

499

229. Schmeier, T. J.; Dobereiner, G. E.; Crabtree, R. H.; Hazari, N. Secondary Coordination Sphere Interactions Facilitate the Insertion Step in an Iridium(III) CO2 Reduction Catalyst. J. Am. Chem. Soc. 2011, 133, 9274–9277. 230. Tai, C.-C.; Chang, T.; Roller, B.; Jessop, P. G. High-Pressure Combinatorial Screening of Homogeneous Catalysts: Hydrogenation of Carbon Dioxide. Inorg. Chem. 2003, 42, 7340–7341. 231. Federsel, C.; Boddien, A.; Jackstell, R.; Jennerjahn, R.; Dyson, P. J.; Scopelliti, R.; Laurenczy, G.; Beller, M. A Well-Defined Iron Catalyst for the Reduction of Bicarbonates and Carbon Dioxide to Formates, Alkyl Formates, and Formamides. Angew. Chem. Int. Ed. 2010, 49, 9777–9780. 232. Langer, R.; Diskin-Posner, Y.; Leitus, G.; Shimon, L. J. W.; Ben-David, Y.; Milstein, D. Low-Pressure Hydrogenation of Carbon Dioxide Catalyzed by an Iron Pincer Complex Exhibiting Noble Metal Activity. Angew. Chem. Int. Ed. 2011, 50, 9948–9952. 233. Khan, M. M. T.; Halligudi, S. B.; Shukla, S. Reduction of CO2 by Molecular Hydrogen to Formic Acid and Formaldehyde and Their Decomposition to CO and H2O. J. Mol. Catal. 1989, 57, 47–60. 234. Miller, A. J. M.; Heinekey, D. M.; Mayer, J. M.; Goldberg, K. I. Catalytic Disproportionation of Formic Acid to Generate Methanol. Angew. Chem. Int. Ed. 2013, 52, 3981–3984. 235. Balaraman, E.; Gunanathan, C.; Zhang, J.; Shimon, L. J. W.; Milstein, D. Efficient Hydrogenation of Organic Carbonates, Carbamates and Formates Indicates Alternative Routes to Methanol Based on CO2 and CO. Nat. Chem. 2011, 3, 609–614. 236. Balaraman, E.; Gnanaprakasam, B.; Shimon, L. J. W.; Milstein, D. Direct Hydrogenation of Amides to Alcohols and Amines Under Mild Conditions. J. Am. Chem. Soc. 2010, 132, 16756–16758. 237. Han, Z.; Rong, L.; Wu, J.; Zhang, L.; Wang, Z.; Ding, K. Catalytic Hydrogenation of Cyclic Carbonates: A Practical Approach From CO2 and Epoxides to Methanol and Diols. Angew. Chem. Int. Ed. 2012, 51, 13041–13045. 238. Waldie, K. M.; Ostericher, A. L.; Reineke, M. H.; Sasayama, A. F.; Kubiak, C. P. Hydricity of Transition-Metal Hydrides: Thermodynamic Considerations for CO2 Reduction. ACS Catal. 2018, 8, 1313–1324. 239. Windle, C. D.; Perutz, R. N. Advances in Molecular Photocatalytic and Electrocatalytic CO2 Reduction. Coord. Chem. Rev. 2012, 256, 2562–2570. 240. Reithmeier, R.; Bruckmeier, C.; Rieger, B. Conversion of CO2 Via Visible Light Promoted Homogeneous Redox Catalysis. Catalysts 2012, 2, 544–571. 241. Boutin, E.; Merakeb, L.; Ma, B.; Boudy, B.; Wang, M.; Bonin, J.; Anxolabéhère-Mallart, E.; Robert, M. Molecular Catalysis of CO2 Reduction: Recent Advances and Perspectives in Electrochemical and Light-Driven Processes With Selected Fe, Ni and Co Aza Macrocyclic and Polypyridine Complexes. Chem. Soc. Rev. 2020, 49, 5772–5809. 242. Manbeck, G. F.; Fujita, E. A Review of Iron and Cobalt Porphyrins, Phthalocyanines and Related Complexes for Electrochemical and Photochemical Reduction of Carbon Dioxide. J. Porphyrins Phthalocyanines 2015, 19, 45–64. 243. Kinzel, N. W.; Werlé, C.; Leitner, W. Transition Metal Complexes as Catalysts for the Electroconversion of CO2: An Organometallic Perspective. Angew. Chem. Int. Ed. 2021https://doi.org/10.1002/anie.202006988. 244. Fisher, B. J.; Eisenberg, R. Electrocatalytic Reduction of Carbon Dioxide by Using Macrocycles of Nickel and Cobalt. J. Am. Chem. Soc. 1980, 102, 7361–7363. 245. Tezuka, M.; Yajima, T.; Tsuchiya, A.; Matsumoto, Y.; Uchida, Y.; Hidai, M. Electroreduction of Carbon Dioxide Catalyzed by Iron-Sulfur Cluster Compounds [Fe4S4(SR)4]2. J. Am. Chem. Soc. 1982, 104, 6834–6836. 246. Hawecker, J.; Lehn, J.-M.; Ziessel, R. Electrocatalytic Reduction of Carbon Dioxide Mediated by Re(bipy)(CO)3Cl (bipy ¼ 2,2[Prime or Minute]-Bipyridine). J. Chem. Soc. Chem. Commun. 1984, ;328–330. 247. Ishida, H.; Tanaka, K.; Tanaka, T. The Electrochemical Reduction of CO2 Catalyzed by Ruthenium Carbonyl Complexes. Chem. Lett. 1985, 14, 405–406. 248. Hammouche, M.; Lexa, D.; Momenteau, M.; Saveant, J. M. Chemical Catalysis of Electrochemical Reactions. Homogeneous Catalysis of the Electrochemical Reduction of Carbon Dioxide by Iron("0") Porphyrins. Role of the Addition of Magnesium Cations. J. Am. Chem. Soc. 1991, 113, 8455–8466. 249. Bhugun, I.; Lexa, D.; Savéant, J.-M. Catalysis of the Electrochemical Reduction of Carbon Dioxide by Iron(0) Porphyrins: Synergystic Effect of Weak Brönsted Acids. J. Am. Chem. Soc. 1996, 118, 1769–1776. 250. Ishida, H.; Fujiki, K.; Ohba, T.; Ohkubo, K.; Tanaka, K.; Terada, T.; Tanaka, T. Ligand Effects of Ruthenium 2,20 -Bipyridine and 1,10-Phenanthroline Complexes on the Electrochemical Reduction of CO2. J. Chem. Soc. Dalton Trans. 1990, ;2155–2160. 251. DuBois, D. L.; Miedaner, A.; Haltiwanger, R. C. Electrochemical Reduction of Carbon Dioxide Catalyzed by [Pd(Triphosphine)(Solvent)](BF4)2 Complexes: Synthetic and Mechanistic Studies. J. Am. Chem. Soc. 1991, 113, 8753–8764. 252. Nagao, H.; Mizukawa, T.; Tanaka, K. Carbon-Carbon Bond Formation in the Electrochemical Reduction of Carbon Dioxide Catalyzed by a Ruthenium Complex. Inorg. Chem. 1994, 33, 3415–3420. 253. Bourrez, M.; Molton, F.; Chardon-Noblat, S.; Deronzier, A. [Mn(bipyridyl)(CO)3Br]: An Abundant Metal Carbonyl Complex as Efficient Electrocatalyst for CO2 Reduction. Angew. Chem. Int. Ed. 2011, 50, 9903–9906. 254. Kang, P.; Cheng, C.; Chen, Z.; Schauer, C. K.; Meyer, T. J.; Brookhart, M. Selective Electrocatalytic Reduction of CO2 to Formate by Water-Stable Iridium Dihydride Pincer Complexes. J. Am. Chem. Soc. 2012, 134, 5500–5503. 255. Taheri, A.; Thompson, E. J.; Fettinger, J. C.; Berben, L. A. An Iron Electrocatalyst for Selective Reduction of CO2 to Formate in Water: Including Thermochemical Insights. ACS Catal. 2015, 5, 7140–7151. 256. Clark, M. L.; Grice, K. A.; Moore, C. E.; Rheingold, A. L.; Kubiak, C. P. Electrocatalytic CO2 Reduction by M(Bpy-R)(CO)4 (M ¼ Mo, W; R ¼ H, tBu) Complexes. Electrochemical, Spectroscopic, and Computational Studies and Comparison With Group 7 Catalysts. Chem. Sci. 2014, 5, 1894–1900. 257. Rosas-Hernández, A.; Junge, H.; Beller, M.; Roemelt, M.; Francke, R. Cyclopentadienone Iron Complexes as Efficient and Selective Catalysts for the Electroreduction of CO2 to CO. Cat. Sci. Technol. 2017, 7, 459–465. 258. Roy, S.; Sharma, B.; Pécaut, J.; Simon, P.; Fontecave, M.; Tran, P. D.; Derat, E.; Artero, V. Molecular Cobalt Complexes With Pendant Amines for Selective Electrocatalytic Reduction of Carbon Dioxide to Formic Acid. J. Am. Chem. Soc. 2017, 139, 3685–3696. 259. Hooe, S. L.; Dressel, J. M.; Dickie, D. A.; Machan, C. W. Highly Efficient Electrocatalytic Reduction of CO2 to CO by a Molecular Chromium Complex. ACS Catal. 2020, 10, 1146–1151. 260. Meshitsuka, S.; Ichikawa, M.; Tamaru, K. Electrocatalysis by Metal Phthalocyanines in the Reduction of Carbon Dioxide. J. Chem. Soc. Chem. Commun. 1974, ;158–159. 261. Ehrenfeld, R. Zur elektrolytischen Reduction der Kohlensäure. Ber. Dtsch. Chem. Ges. 1905, 38, 4138–4143. 262. Fischer, F.; Prziza, O. Über die elektrolytische Reduktion von unter Druck gelöstem Kohlendioxyd und Kohlenoxyd. Ber. Dtsch. Chem. Ges. 1914, 47, 256–260. 263. Royer, M. E. Reduction of Carbonic Acid Into Formic Acid. CR Acad. Sci. Paris 1870, 70, 73. 264. Coehn, A.; Jahn, S. Ueber elektrolytische Reduction der Kohlensäure. Ber. Dtsch. Chem. Ges. 1904, 37, 2836–2842. 265. Rabinowitsch, M.; Maschowetz, A. Elektrochemische Gewinnung von Formiaten aus Kohlensäure. Z. Elektrochem. Angew. Phys. Chem. 1930, 36, 846–850. 266. Teeter, T. E.; Van Rysselberghe, P. Reduction of Carbon Dioxide on Mercury Cathodes. J. Chem. Phys. 1954, 22, 759–760. 267. Rysselberghe, P. V.; Alkire, G. J. Polarographic Reduction of Carbon Dioxide. J. Am. Chem. Soc. 1944, 66, 1801. 268. Rysselberghe, P. V. Polarographic Reduction of Carbon Dioxide. II. Polymerization and Adsorption at the Dropping Mercury Cathode1. J. Am. Chem. Soc. 1946, 68, 2047–2049. 269. Rysselberghe, P.v.; Alkire, G. J.; McGee, J. M. Polarographic Reduction of Carbon Dioxide. III. Description and Interpretation of the Waves1. J. Am. Chem. Soc. 1946, 68, 2050–2055. 270. Paik, W.; Andersen, T. N.; Eyring, H. Kinetic Studies of the Electrolytic Reduction of Carbon Dioxide on the Mercury Electrode. Electrochim. Acta 1969, 14, 1217–1232. 271. Hiratsuka, K.; Takahashi, K.; Sasaki, H.; Toshima, S. Electrocatalytic Behavior of Tetrasulfonated Metal Phthalocyanines in the Reduction of Carbon Dioxide. Chem. Lett. 1977, 6, 1137–1140. 272. Takahashi, K.; Hiratsuka, K.; Sasaki, H.; Toshima, S. Electrocatalytic Behavior of Metal Porphyrins in the Reduction of Carbon Dioxide. Chem. Lett. 1979, 8, 305–308.

500

Introduction to the Organometallic Chemistry of Carbon Dioxide

273. Beley, M.; Collin, J.-P.; Ruppert, R.; Sauvage, J.-P. Nickel(II)-Cyclam: An Extremely Selective Electrocatalyst for Reduction of CO2 in Water. J. Chem. Soc. Chem. Commun. 1984, ;1315–1316. 274. Feroci, G.; Roffia, S. On the Reduction of Oxygen in Dimethylformamide. J. Electroanal. Chem. Interfacial Electrochem. 1976, 71, 191–198. 275. Hawecker, J.; Lehn, J.-M.; Ziessel, R. Efficient Photochemical Reduction of CO2 to CO by Visible Light Irradiation of Systems Containing Re(Bipy)(CO)3X or Ru(Bipy)32 +–Co2 + Combinations as Homogeneous Catalysts. J. Chem. Soc. Chem. Commun. 1983, ;536–538. 276. Slater, S.; Wagenknecht, J. H. Electrochemical Reduction of Carbon Dioxide Catalyzed by Rh(Diphos)2Cl. J. Am. Chem. Soc. 1984, 106, 5367–5368. 277. Ishida, H.; Tanaka, H.; Tanaka, K.; Tanaka, T. Selective Formation of HCOO- in the Electrochemical CO2 Reduction Catalysed by [Ru(Bpy)2(CO)2]2+(Bpy ¼ 2,2[Prime or Minute]-Bipyridine). J. Chem. Soc. Chem. Commun. 1987, ;131–132. 278. Bruce, M. R. M.; Megehee, E.; Sullivan, B. P.; Thorp, H. H.; O’Toole, T. R.; Downard, A.; Pugh, J. R.; Meyer, T. J. Electrocatalytic Reduction of Carbon Dioxide Based on 2,20 -Bipyridyl Complexes of Osmium. Inorg. Chem. 1992, 31, 4864–4873. 279. Hammouche, M.; Lexa, D.; Savéant, J. M.; Momenteau, M. Catalysis of the Electrochemical Reduction of Carbon Dioxide by Iron(“0”) Porphyrins. J. Electroanal. Chem. Interfacial Electrochem. 1988, 249, 347–351. 280. Costentin, C.; Drouet, S.; Robert, M.; Savéant, J.-M. A Local Proton Source Enhances CO2 Electroreduction to CO by a Molecular Fe Catalyst. Science 2012, 338, 90–94. 281. Costentin, C.; Passard, G.; Robert, M.; Savéant, J.-M. Pendant Acid–Base Groups in Molecular Catalysts: H-Bond Promoters or Proton Relays? Mechanisms of the Conversion of CO2 to CO by Electrogenerated Iron(0)Porphyrins Bearing Prepositioned Phenol Functionalities. J. Am. Chem. Soc. 2014, 136, 11821–11829. 282. Azcarate, I.; Costentin, C.; Robert, M.; Savéant, J.-M. Through-Space Charge Interaction Substituent Effects in Molecular Catalysis Leading to the Design of the Most Efficient Catalyst of CO2-to-CO Electrochemical Conversion. J. Am. Chem. Soc. 2016, 138, 16639–16644. 283. Azcarate, I.; Costentin, C.; Robert, M.; Savéant, J.-M. Dissection of Electronic Substituent Effects in Multielectron–Multistep Molecular Catalysis. Electrochemical CO2-to-CO Conversion Catalyzed by Iron Porphyrins. J. Phys. Chem. C 2016, 120, 28951–28960. 284. Nichols, E. M.; Derrick, J. S.; Nistanaki, S. K.; Smith, P. T.; Chang, C. J. Positional Effects of Second-Sphere Amide Pendants on Electrochemical CO2 Reduction Catalyzed by Iron Porphyrins. Chem. Sci. 2018, 9, 2952–2960. 285. Gotico, P.; Boitrel, B.; Guillot, R.; Sircoglou, M.; Quaranta, A.; Halime, Z.; Leibl, W.; Aukauloo, A. Second-Sphere Biomimetic Multipoint Hydrogen-Bonding Patterns to Boost CO2 Reduction of Iron Porphyrins. Angew. Chem. Int. Ed. 2019, 58, 4504–4509. 286. Chardon-Noblat, S.; Deronzier, A.; Ziessel, R.; Zsoldos, D. Selective Synthesis and Electrochemical Behavior of Trans(cl)- and Cis(cl)-[Ru(Bpy)(CO)2Cl2] Complexes (Bpy ¼ 2,2‘Bipyridine). Comparative Studies of their Electrocatalytic Activity toward the Reduction of Carbon Dioxide. Inorg. Chem. 1997, 36, 5384–5389. 287. Collomb-Dunand-Sauthier, M. N.; Deronzier, A.; Ziessel, R. Electrocatalytic Reduction of Carbon-Dioxide With Mono(Bipyridine)Carbonylruthenium Complexes in Solution or as Polymeric Thin-Films. Inorg. Chem. 1994, 33, 2961–2967. 288. Chardon-Noblat, S.; Deronzier, A.; Hartl, F.; van Slageren, J.; Mahabiersing, T. A Novel Organometallic Polymer of Osmium(0), [Os(2,20 -Bipyridine)(CO)2]N: Its Electrosynthesis and Electrocatalytic Properties Towards CO2 Reduction. Eur. J. Inorg. Chem. 2001, 2001, 613–617. 289. Chen, Z.; Chen, C.; Weinberg, D. R.; Kang, P.; Concepcion, J. J.; Harrison, D. P.; Brookhart, M. S.; Meyer, T. J. Electrocatalytic Reduction of CO2 to CO by Polypyridyl Ruthenium Complexes. Chem. Commun. 2011, 47, 12607–12609. 290. Gonell, S.; Massey, M. D.; Moseley, I. P.; Schauer, C. K.; Muckerman, J. T.; Miller, A. J. M. The Trans Effect in Electrocatalytic CO2 Reduction: Mechanistic Studies of Asymmetric Ruthenium Pyridyl-Carbene Catalysts. J. Am. Chem. Soc. 2019, 141, 6658–6671. 291. Gonell, S.; Lloret-Fillol, J.; Miller, A. J. M. An Iron Pyridyl-Carbene Electrocatalyst for Low Overpotential CO2 Reduction to CO. ACS Catal. 2020, ;615–626. 292. Gonell, S.; Assaf, E. A.; Duffee, K. D.; Schauer, C. K.; Miller, A. J. M. Kinetics of the Trans Effect in Ruthenium Complexes Provide Insight Into the Factors That Control Activity and Stability in CO2 Electroreduction. J. Am. Chem. Soc. 2020, 142, 8980–8999. 293. Derrick, J. S.; Loipersberger, M.; Chatterjee, R.; Iovan, D. A.; Smith, P. T.; Chakarawet, K.; Yano, J.; Long, J. R.; Head-Gordon, M.; Chang, C. J. Metal–Ligand Cooperativity Via Exchange Coupling Promotes Iron- Catalyzed Electrochemical CO2 Reduction at Low Overpotentials. J. Am. Chem. Soc. 2020, 142, 20489–20501. 294. Grills, D. C.; Farrington, J. A.; Layne, B. H.; Lymar, S. V.; Mello, B. A.; Preses, J. M.; Wishart, J. F. Mechanism of the Formation of a Mn-Based CO2 Reduction Catalyst Revealed by Pulse Radiolysis With Time-Resolved Infrared Detection. J. Am. Chem. Soc. 2014, 136, 5563–5566. 295. Sampson, M. D.; Nguyen, A. D.; Grice, K. A.; Moore, C. E.; Rheingold, A. L.; Kubiak, C. P. Manganese Catalysts With Bulky Bipyridine Ligands for the Electrocatalytic Reduction of Carbon Dioxide: Eliminating Dimerization and Altering Catalysis. J. Am. Chem. Soc. 2014, 136, 5460–5471. 296. Sampson, M. D.; Kubiak, C. P. Manganese Electrocatalysts With Bulky Bipyridine Ligands: Utilizing Lewis Acids to Promote Carbon Dioxide Reduction at Low Overpotentials. J. Am. Chem. Soc. 2016, 138, 1386–1393. 297. Machan, C. W.; Sampson, M. D.; Chabolla, S. A.; Dang, T.; Kubiak, C. P. Developing a Mechanistic Understanding of Molecular Electrocatalysts for CO2 Reduction Using Infrared Spectroelectrochemistry. Organometallics 2014, 33, 4550–4559. 298. Ngo, K. T.; McKinnon, M.; Mahanti, B.; Narayanan, R.; Grills, D. C.; Ertem, M. Z.; Rochford, J. Turning on the Protonation-First Pathway for Electrocatalytic CO2 Reduction by Manganese Bipyridyl Tricarbonyl Complexes. J. Am. Chem. Soc. 2017, 139, 2604–2618. 299. Rail, M. D.; Berben, L. A. Directing the Reactivity of [HFe4N(CO)12]- Toward H+ or CO2 Reduction by Understanding the Electrocatalytic Mechanism. J. Am. Chem. Soc. 2011, 133, 18577–18579. 300. Loewen, N. D.; Neelakantan, T. V.; Berben, L. A. Renewable Formate from C–H Bond Formation With CO2: Using Iron Carbonyl Clusters as Electrocatalysts. Acc. Chem. Res. 2017, 50, 2362–2370. 301. Kanega, R.; Ertem, M. Z.; Onishi, N.; Szalda, D. J.; Fujita, E.; Himeda, Y. CO2 Hydrogenation and Formic Acid Dehydrogenation Using Ir Catalysts With Amide-Based Ligands. Organometallics 2020, 39, 1519–1531. 302. Grice, K. A. Carbon Dioxide Reduction With Homogenous Early Transition Metal Complexes: Opportunities and Challenges for Developing CO2 Catalysis. Coord. Chem. Rev. 2017, 336, 78–95. 303. Jiang, C.; Nichols, A. W.; Machan, C. W. A Look at Periodic Trends in d-Block Molecular Electrocatalysts for CO2 Reduction. Dalton Trans. 2019, 48, 9454–9468. 304. Moreno, J. J.; Hooe, S. L.; Machan, C. W. A DFT Study on the Electrocatalytic Reduction of CO2 to CO by a Molecular Chromium Complex. Inorg. Chem. 2021https://doi.org/ 10.1021/acs.inorgchem.1020c03136. 305. Nichols, A. W.; Chatterjee, S.; Sabat, M.; Machan, C. W. Electrocatalytic Reduction of CO2 to Formate by an Iron Schiff Base Complex. Inorg. Chem. 2018, 57, 2111–2121. 306. Nichols, A. W.; Hooe, S. L.; Kuehner, J. S.; Dickie, D. A.; Machan, C. W. Electrocatalytic CO2 Reduction to Formate With Molecular Fe(III) Complexes Containing Pendent Proton Relays. Inorg. Chem. 2020, 59, 5854–5864. 307. Aresta, M.; Dibenedetto, A.; Angelini, A. The Use of Solar Energy Can Enhance the Conversion of Carbon Dioxide Into Energy-Rich Products: Stepping towards Artificial Photosynthesis. Phil. Trans. R. Soc. A 2013, 371https://doi.org/10.1098/rsta.2012.0111. 308. Derien, S.; Dunach, E.; Perichon, J. From Stoichiometry to Catalysis: Electroreductive Coupling of Alkynes and Carbon Dioxide with Nickel-Bipyridine Complexes. Magnesium Ions as the Key for Catalysis. J. Am. Chem. Soc. 1991, 113, 8447–8454. 309. Nédélec, J.-Y.; Périchon, J.; Troupel, M. Organic Electroreductive Coupling Reactions Using Transition Metal Complexes as Catalysts. In Electrochemistry VI Electroorganic Synthesis: Bond Formation at Anode and Cathode; Steckhan, E., Ed.; Topics in Current Chemistry Springer Berlin: Heidelberg, 1997; vol. 185; pp 141–173. 310. Labbé, E.; Duñach, E.; Périchon, J. Ligand-Directed Reaction Products in the Nickel-Catalyzed Electrochemical Carboxylation of Terminal Alkynes. J. Organomet. Chem. 1988, 353, C51–C56. 311. Jin, D.; Schmeier, T. J.; Williard, P. G.; Hazari, N.; Bernskoetter, W. H. Lewis Acid Induced b-Elimination from a Nickelalactone: Efforts toward Acrylate Production From CO2 and Ethylene. Organometallics 2013, 32, 2152–2159.

Introduction to the Organometallic Chemistry of Carbon Dioxide

501

312. Takimoto, M.; Nakamura, Y.; Kimura, K.; Mori, M. Highly Enantioselective Catalytic Carbon Dioxide Incorporation Reaction: Nickel-Catalyzed Asymmetric Carboxylative Cyclization of Bis-1,3-Dienes. J. Am. Chem. Soc. 2004, 126, 5956–5957. 313. Williams, C. M.; Johnson, J. B.; Rovis, T. Nickel-Catalyzed Reductive Carboxylation of Styrenes Using CO2. J. Am. Chem. Soc. 2008, 130, 14936–14937. 314. Fujihara, T.; Nogi, K.; Xu, T.; Terao, J.; Tsuji, Y. Nickel-Catalyzed Carboxylation of Aryl and Vinyl Chlorides Employing Carbon Dioxide. J. Am. Chem. Soc. 2012, 134, 9106–9109. 315. León, T.; Correa, A.; Martin, R. Ni-Catalyzed Direct Carboxylation of Benzyl Halides With CO2. J. Am. Chem. Soc. 2013, 135, 1221–1224. 316. Yoo, C.; Oh, S.; Kim, J.; Lee, Y. Transmethylation of a Four-Coordinate Nickel(I) Monocarbonyl Species With Methyl Iodide. Chem. Sci. 2014, 5, 3853–3858. 317. Ang, N. W. J.; Oliveira, J. C. A.; Ackermann, L. Electroreductive Cobalt-Catalyzed Carboxylation: Cross-Electrophile Electrocoupling With Atmospheric CO2. Angew. Chem. Int. Ed. 2020, 59, 12842–12847. 318. Inoue, S.; Koinuma, H.; Tsuruta, T. Copolymerization of Carbon Dioxide and Epoxide. J. Polym. Sci., Part B: Polym. Lett. 1969, 7, 287–292. 319. Inoue, S.; Koinuma, H.; Tsuruta, T. Copolymerization of Carbon Dioxide and Epoxide With Organometallic Compounds. Makromol. Chem. 1969, 130, 210–220. 320. Inoue, S. Immortal Polymerization: The Outset, Development, and Application. J. Polym. Sci. A Polym. Chem. 2000, 38, 2861–2871. 321. Takeda, N.; Inoue, S. Polymerization of 1,2-Epoxypropane and Copolymerization With Carbon Dioxide Catalyzed by Metalloporphyrins. Die Makromolekulare Chemie 1978, 179, 1377–1381. 322. Kruper, W. J.; Dellar, D. D. Catalytic Formation of Cyclic Carbonates From Epoxides and CO2 With Chromium Metalloporphyrinates. J. Org. Chem. 1995, 60, 725–727. 323. Darensbourg, D. J.; Holtcamp, M. W.; et al. Macromolecules 1995, 28, 7577–7579. 324. Darensbourg, D. J.; Holtcamp, M. W.; Struck, G. E.; Zimmer, M. S.; Niezgoda, S. A.; Rainey, P.; Robertson, J. B.; Draper, J. D.; Reibenspies, J. H. Catalytic Activity of a Series of Zn(II) Phenoxides for the Copolymerization of Epoxides and Carbon Dioxide. J. Am. Chem. Soc. 1999, 121, 107–116. 325. Cheng, M.; Lobkovsky, E. B.; Coates, G. W. Catalytic Reactions Involving C1 Feedstocks: New High-Activity Zn(II)-Based Catalysts for the Alternating Copolymerization of Carbon Dioxide and Epoxides. J. Am. Chem. Soc. 1998, 120, 11018–11019. 326. Sugimoto, H.; Ohshima, H.; Inoue, S. Alternating Copolymerization of Carbon Dioxide and Epoxide by Manganese Porphyrin: The First Example of Polycarbonate Synthesis from 1-Atm Carbon Dioxide. J. Polym. Sci. A Polym. Chem. 2003, 41, 3549–3555. 327. Jacobsen, E. N.; Tokunaga, M.; Larrow, J. F. PCT Int. Appl. WO 00/09463. 328. Paddock, R. L.; Nguyen, S. T. Chemical CO2 Fixation: Cr(III) Salen Complexes As Highly Efficient Catalysts for the Coupling of CO2 and Epoxides. J. Am. Chem. Soc. 2001, 123, 11498–11499. 329. Darensbourg, D. J.; Yarbrough, J. C. Mechanistic Aspects of the Copolymerization Reaction of Carbon Dioxide and Epoxides, Using a Chiral Salen Chromium Chloride Catalyst. J. Am. Chem. Soc. 2002, 124, 6335–6342. 330. Qin, Z.; Thomas, C. M.; Lee, S.; Coates, G. W. Cobalt-Based Complexes for the Copolymerization of Propylene Oxide and CO2: Active and Selective Catalysts for Polycarbonate Synthesis. Angew. Chem. Int. Ed. 2003, 42, 5484–5487. 331. Darensbourg, D. J.; Billodeaux, D. R. Aluminum Salen Complexes and Tetrabutylammonium Salts: A Binary Catalytic System for Production of Polycarbonates From CO2 and Cyclohexene Oxide. Inorg. Chem. 2005, 44, 1433–1442. 332. Darensbourg, D. J.; Adams, M. J.; Yarbrough, J. C. Toward the Design of Double Metal Cyanides for the Copolymerization of CO2 and Epoxides. Inorg. Chem. 2001, 40, 6543–6544. 333. Darensbourg, D. J.; Adams, M. J.; Yarbrough, J. C.; Phelps, A. L. Synthesis and Structural Characterization of Double Metal Cyanides of Iron and Zinc: Catalyst Precursors for the Copolymerization of Carbon Dioxide and Epoxides. Inorg. Chem. 2003, 42, 7809–7818. 334. Buchard, A.; Kember, M. R.; Sandeman, K. G.; Williams, C. K. A Bimetallic Iron(III) Catalyst for CO2/Epoxide Coupling. Chem. Commun. 2011, 47, 212–214. 335. Nakano, K.; Kobayashi, K.; Ohkawara, T.; Imoto, H.; Nozaki, K. Copolymerization of Epoxides With Carbon Dioxide Catalyzed by Iron–Corrole Complexes: Synthesis of a Crystalline Copolymer. J. Am. Chem. Soc. 2013, 135, 8456–8459. 336. Burgess, S. A.; Appel, A. M.; Linehan, J. C.; Wiedner, E. S. Changing the Mechanism for CO2 Hydrogenation Using Solvent-Dependent Thermodynamics. Angew. Chem. Int. Ed. 2017, 56, 15002–15005. 337. Assen, N. V. D.; Sternberg, A.; Kätelhön, A.; Bardow, A. Environmental Potential of Carbon Dioxide Utilization in the Polyurethane Supply Chain. Faraday Discuss. 2015, 183, 291–307. 338. Barzagli, F.; Mani, F.; Peruzzini, M. From Greenhouse Gas to Feedstock: Formation of Ammonium Carbamate From CO2 and NH3 in Organic Solvents and Its Catalytic Conversion Into Urea Under Mild Conditions. Green Chem. 2011, 13, 1267–1274.

1.16

Alkane s-Complexes

Rowan D Young, National University of Singapore, Singapore, Singapore © 2022 Elsevier Ltd. All rights reserved.

1.16.1 1.16.1.1 1.16.1.2 1.16.2 1.16.2.1 1.16.2.1.1 1.16.2.1.2 1.16.2.2 1.16.2.2.1 1.16.2.2.2 1.16.2.3 1.16.3 1.16.3.1 1.16.3.2 1.16.3.3 1.16.3.4 1.16.4 References

Introduction Alkane s-complexes in context Synthesis of alkane s-complexes Alkane s-complex characterization Solution studies of alkane s-complexes Fast-spectroscopic observations NMR spectroscopically observed alkane s-complexes Structurally characterized alkane s-complexes Single crystal X-ray diffraction (SCXRD) Neutron diffraction studies Alkane complexes studied computationally Bonding trends in alkane s-complexes Bonding modes Bonding distance and angles Alkane complexes versus fluxional alkyl hydrides Binding selectivity and dynamic exchange processes Summary

1.16.1

Introduction

1.16.1.1

Alkane s-complexes in context

502 502 503 504 504 504 505 508 508 510 512 513 513 514 515 516 518 519

Alkane s-complexes belong to the larger class of metal complexes known as s-complexes.1 s-Complexes are so called not due to the symmetry of the bond formed between the ligand and the metal, but rather due to the origin of the electrons involved in donation to the metal center. In s-complexes, electrons in energetically stable s-bonding molecular orbitals in one or more ligands donate to the metal center. Somewhat confusingly, the primary bond formed from the donation of electrons from the ligand’s s-bond to a vacant metal orbital is a s-bond (i.e. it has s-symmetry). Given the similarity in the terms used to describe different aspects of this bonding, it is essential to be specific when describing such interactions and complexes. In an analogous bonding scenario, alkene ligands bind to metals using electrons housed in p-symmetry carbon-carbon bonds to form s-bonds with the metal.2 Of course, as outlined by the Dewar-Chatt-Duncanson model,3 alkene ligands can act as electron acceptors from suitably aligned occupied metal d-orbitals to form additional metal-ligand interactions of p-symmetry. Thus, these ligands are termed p-acidic (i.e. p-symmetry electron acceptors). The bonding in ligands that bind through s-bond electrons (s-bound ligands) can also be described using the Dewar-Chatt-Duncanson model, where electron retrodonation from the metal to the ligand involves off-axis d orbitals of the metal and s anti-bonding orbitals on the ligand (i.e. s-bound ligands are also p-acidic). Although this interaction involves anti-bonding orbitals housed on the ligand, it has metal-ligand bonding character, and leads to the formation of a p-symmetry bond between the metal and ligand (Fig. 1). Electronically, these two interactions can describe much of the chemistry observed in all s-complexes. For example, depopulation of a s-bonding orbital in a s-coordinated ligand through donation of electrons from this orbital to a metal center, combined with population of a s anti-bonding orbital on the ligand through metal retrodonation, reduces the bond order of the coordinated bond, leading to bond cleavage and oxidative addition of the bond to the metal. As such, metals are capable of binding very weakly donating s-bound ligands, and cleaving very strong element-element bonds. This is highlighted in the case of the dihydrogen ligand, which possesses only two electrons housed in a very stable s-bonding orbital. In scenarios where an ‘intact’ dihydrogen molecule coordinates to a metal center, electrons housed in the hydrogen’s s-bonding orbital are the only electron source for dative bond formation. And the detrimental effect that p-retrodonation has to the H–H bond strength can also explain how transition metals are able to cleave relatively strong molecular hydrogen bonds. There are a multitude of examples where hydrogen is able to (1) act as a stable ligand,4 and (2) be cleaved into two hydride ligands5 through the synergistic s-donation/p-retrodonation interaction with metal centers. Apart from dihydrogen, many stable s-bound ligands are known that bind through group 13 and 14 element-hydrogen bonds. Full characterization of B–H, Al–H, Si–H, Ge–H and Sn–H s-bound ligands is well documented.6 In these instances, the ability of a metal to retrodonate to the s antibonding orbitals is sterically compromised by the substituents on the group 13/14 element (as compared to the dihydrogen model), however, the more highly polarized Ed+–Hd− bonds provide a stronger s-donor ability for these ligands (as compared to dihydrogen). Notably, missing from the list of stable s-bound ligands whose complexes are able to be fully characterized are alkane ligands that bind through a C–H bond.7C–H bonds in alkane ligands suffer from steric hindrance as

502

Comprehensive Organometallic Chemistry IV

https://doi.org/10.1016/B978-0-12-820206-7.00001-9

Alkane s-Complexes

503

Fig. 1 (Top) Electronic interaction between a C–H s-bond and a rhenium metal center as described by the Dewar-Chatt-Duncanson model. (Bottom) Alkane scomplex as an intermediate of C–H activation.

with other E–H bound ligands (E 6¼ H), but the C–H bond is much less polar, and far more covalent in nature than other E–H bonds, resulting in a very poor s-donation ability. The combination of these factors makes alkane ligands intrinsically unstable, and to date, there exist no examples of alkane complexes that can be characterized in both solution and the solid-state. Despite this, a great deal is known regarding these complexes from (1) partial characterization of judiciously designed (or serendipitously discovered) metal systems that offer greater alkane s-complex stability, (2) comparison to other well-characterized s-complexes, and (3) comparison to complexes containing agostic interactions (chelate tethered C–H s-bound ligands). The precarious local energy minimum that alkane s-complexes occupy as intermediates en route to alkane transition metal mediated C–H activation renders such complexes of extreme interest to contemporary chemists (Fig. 1).8 Given that the primary organic feedstock of fossil fuel is alkanes, mild methods of controlled functionalization are highly sought. Our improved understanding of alkane coordination and activation has already led to a number of methods for the functionalization of unactivated alkanes, including methane.9 This chapter summarizes known alkane complexes and their characterization data, and outlines our understanding of alkane binding to metals based on these data.

1.16.1.2

Synthesis of alkane s-complexes

Given the weak binding nature of alkane ligands, traditional methods of ligand substitution cannot be used to access alkane s-complexes. Indeed, even renowned inert solvents and weakly coordinating anions (WCA) preferentially coordinate to metal centers in place of alkane ligands. Often alkane s-complexes represent kinetic products, thus strategies to access and characterize these complexes often rely upon the use of low temperatures and sterically encumbered weakly coordinating anions. Further, the avoidance of electron donating solvents and anions can prevent the formation of thermodynamically preferred solvent/anion adducts. To date, synthetic strategies to access alkane s-complexes can be categorized into four groups. Namely; (1) coordination of an alkane to a stable, otherwise unsaturated metal center (Fig. 2A), (2) photo-dissociation of a ligand to provide a vacant coordination site for coordination of an alkane (Fig. 2B), (3) protonation of a metal alkyl complex (Fig. 2C), and (4) hydrogenation of an alkene ligand to generate an alkane ligand in the coordination sphere of a metal (Fig. 2D). Coordination of an alkane directly to a coordinatively unsaturated metal center rarely results in observable alkane complexes, given that alkane ligands offer little extra stability to existing stable complexes (Fig. 2A). However, such intermediates are of great interest given their importance to many C–H functionalization reactions.8 Despite the inconsistency of this approach, solid state structures of suspected alkane s-complexes have been reported from the crystallization of metal complexes in alkane solvent systems.10 In these cases, significant stabilization can be offered to the alkane ligand through the creation of a lipophilic pocket, and these interactions may dominate the binding interaction of the alkane to the metal complex.

504

Alkane s-Complexes

(B)

(A)

(C)

(D)

Fig. 2 The various reported methods of synthesizing alkane complexes, including: (A) coordination to an unsaturated metal supported by a ligand cavity architecture; (B) photo-induced loss of a labile ligand to generate a reactive site to which an alkane molecule can coordinate; (C) protonation of an alkyl ligand; (D) hydrogenation of an alkene ligand to produce a coordinated alkane.

An extension of coordination of an alkane to an unsaturated complex is the photogeneration of an unstable coordinatively unsaturated 16 (or less) electron metal center formed by the ejection of one or more ligands, which quickly ligates an alkane solvent molecule (Fig. 2B).11 This method was used to create the first observed alkane s-complex,12 and was used to produce the first alkane oxidative addition products, where an alkane s-complex intermediate would have been expected.13 Precursor complexes for this type of alkane s-complex production necessitate the incorporation of stable leaving groups that are able to dissociate from the metal center. Most commonly, carbon monoxide is the ligand of choice, however, dinitrogen14 and dihydride13a groups have also been used to fulfil this purpose. The production of alkane complexes photolytically has many advantages and is thus the most common method employed in reports of alkane s-complexes. The exact time when complexes are created can be controlled with a light source. Indeed, laser flash photolysis (LFP) allows near complete conversion of starting material to product almost instantly, simplifying the kinetic studies of alkane complexes generated in this manner. The controlled formation of the reactive species in situ is also feasible, as light sources can be engineered to deliver radiation directly to the sample under study.15 Alkane complexes have also been generated by the protonation of metal alkyl species (Fig. 2C).16 This approach is reliant upon the alkane s-complex being more stable than the corresponding activation product (i.e. the alkyl hydrido complex). Protonation may occur directly at the alkyl carbon position or at the metal center followed by fast reductive elimination. The alkane s-complexes tend to be kinetic products, which decompose via displacement of the alkane ligands with solvent molecule or counter anion donors. As such, observation of these complexes has been limited to low temperature NMR experiments where protonation of a metal alkyl to generate an alkane s-complex is carried out on precooled samples. Most recently, alkane s-complexes generated from the hydrogenation of precursor metal olefin complexes have been characterized in the solid state (Fig. 2D).7d,17,18 Alkane s-complexes are frequently generated as intermediates in alkene hydrogenation, but the resulting alkane ligands are readily displaced by solvent, counter anion or more alkene. However, when bound alkenes are hydrogenated in the solid state, it has been found that suitable anions form protective cages that stabilize the metal-alkane interaction, in some cases, even at room temperature for hours. The kinetic stability of such complexes has allowed the use of alkane s-complexes in synthetic applications.17g–j Although a range of approaches have been established to generate observable alkane s-complexes, and collective data of alkane s-complexes reported are ever-growing, no alkane s-complex has yet been characterized in both the solid state and in solution. This represents a huge shortcoming in our ability to experimentally compare alkane ligand binding modes and bond strengths obtained via solution or solid-state characterization methods. Direct comparisons are often made out of necessity between solid state and solution alkane s-complexes, even though the degree of interaction as determined by calculated binding energies can be drastically different in each case. Thus, it can be helpful to examine alkane s-complexes based on the method of observation rather than based on the common chemical features of any system.

1.16.2

Alkane s-complex characterization

1.16.2.1

Solution studies of alkane s-complexes

1.16.2.1.1

Fast-spectroscopic observations

Fast spectroscopic methods that allow measurements on the picosecond scale (e.g. UV–Vis, FTIR spectroscopies) provided the first data for alkane s-complexes and have been utilized to obtain kinetic data for such complexes. However, such methods provide very

Alkane s-Complexes

505

limited information on the bonding interaction between the alkane and metal center, and require the inclusion of a chromophore or a spectator ligand able to provide spectroscopic feedback on the state of the metal complex. Nonetheless, time-resolved IR and matrix isolation UV-Vis and FTIR studies have played a pivotal role in identifying alkane s-complexes. The seminal evidence for alkane s-complexes was obtained by Stolz et al. in 1962.12a Photolysis of group 6 carbonyls at 77 K in isopentane/methylcyclohexane glasses generated species of the type [M(CO)5(alkane)] (M ¼ Cr, Mo, W). At the time, the concept of s-complexes had not even been developed, and the spectroscopic data were understandably interpreted in terms of a free 16-electron ‘M(CO)5’ species in an alkane matrix, rather than an alkane s-complex. Turner subsequently proposed that the pentacarbonyl species were forming weak bonds with the matrix support molecules, but the nature of the interaction was never discussed.12b As a greater understanding of related s-complexes was developed, spectroscopic data were interpreted in terms of alkane C–H bonds acting as L type ligands. Additionally, time-resolved IR spectroscopy was key in providing the first substantial evidence that alkane s-complexes existed as short lived intermediates prior to C–H activation of various alkanes. Studies by Bergman and Harris with photolytically generated ‘Tp Rh(CO)’ {Tp ¼ hydrotris(3,5-dimethyl-pyrazol-1-yl)borate} species showed that alkanes coordinated the rhodium metal center prior to formation of alkyl hydride products, with the pre-activation complexes (i.e. alkane complexes) displaying lifetimes of only a few hundred nanoseconds (Fig. 3).19 Owing to the valuable kinetic data that time-resolved IR can provide at ambient temperatures and the ease of potential alkane complex screening, IR is now more frequently used in combination with NMR to characterize alkane complexes with higher degrees of confidence than as a stand-alone technique. More recently, George undertook a combined time-resolved IR/XAFS study to probe n-heptane binding to the ‘W(CO)5’ fragment.20 XAFS is sensitive to small structural changes, and was able to provide an averaged metal-carbon bond distance for the alkane ligand. In combination with DFT analysis, this approach was able to provide very basic structural information on an alkane complex, [W(CO)5(n-heptane)]. IR spectroscopy still represents the characterization method with the most comprehensive observational data of alkane complexes, and the only method to provide experimental data for vanadium, niobium, tantalum, chromium, molybdenum and ruthenium alkane complexes, as can be seen in Table 1.

1.16.2.1.2

NMR spectroscopically observed alkane s-complexes

There have existed experimental data for alkane s-complexes since 1962,12a however, none of these data could verify the structural composition of such complexes. In particular, metrics for metal-alkane ligand bond distances and angles were non-existent. Models for the bonding in alkane s-complexes were based upon other related s-complexes (e.g. s-silane or s-borane complexes) or agostic interactions. However, in 1998 Ball provided the first NMR evidence for the existence of alkane s-complexes, ‘shedding light’ onto the bonding characteristics and dynamics of alkane ligands coordinated to transition metal centres.39 In contrast to the fast spectroscopic techniques discussed above, NMR spectroscopy offers the ability to observe coordinated alkane carbon and hydrogen atoms directly and, in some cases, to observe the metal center indirectly also. NMR techniques offer information on bond strength, atomic distances, bonding sites, and kinetic data on exchange between bonding sites. Unfortunately, NMR observations require measurement times on the order of a second, much longer than the lifetimes of known solution alkane s-complexes at room temperature. Following on from a report on the relative stability of [CpRe(CO)2(n-C7H16)],28e Ball realized that such a complex may be sufficiently long-lived at low temperature for NMR observation, and he subsequently reported NMR data for the photogenerated complex [CpRe(CO)2(cycloC5H10)] below 193 K (Fig. 4).39 At low temperatures, the lifetime of relatively stable solution alkane s-complexes can be extended from milliseconds (at room temperature) to minutes, suitable for NMR observation. Importantly, the NMR data obtained correlated well with IR observations previously reported, and subsequent NMR/IR combined analyses of photogenerated alkane complexes have been reported, verifying the veracity of alkane s-complexes reported on the strength of IR spectroscopic data. The pioneering work of Ball allowed observation of the weakening of the coordinated C–H bonds (from retrodonation to the C–H s orbital) through measurements of C–H coupling constants (1JCH) (Fig. 4). Additionally, this work allowed them to distinguish specific C–H binding sites. In contrast to what is observed in transition metal mediated C–H activation, methine and

Fig. 3 Laser flash photolysis (LFP) of [Tp Rh(CO)2] in cyclohexane solution ejects a carbonyl ligand and leads to a short-lived alkane complex prior to C–H activation that has been observed using TRIR spectroscopy.19

506

Alkane s-Complexes

Table 1

Alkane s-complexes [LnM(alkane)] detected by infrared spectroscopy.

‘MLn’ Group 5 (Z5-C5H5)M(CO)3 (M ¼ V, Nb, Ta) Group 6 M(CO)5 (M ¼ Cr, Mo, W) (Z5-C5H5)W(Et)(CO)2 (Z6-C6R6)M(CO)2 (R ¼ H, Me, Et; M ¼ Cr, W) (Z5-C5Me5)W(Bpin)(CO)2 fac-(dfepe)Cr(CO)3 Group 7 RMn(CO)4 (R ¼ H, Me) Mn(dmpe)2H Re(dmpe)H (Z5-C5R5)Re(CO)(L) (R ¼ H, Me, iPr, tBu, Ph; L ¼ CO, PF3) (Z5-C5H5)Mn(CO)2 (Z5-C5Me5)Re(Et)2(CO) Tp’Re(CO)2 (Tp’ ¼ Tp, Tp ) KpRe(CO)2 Group 8 Fe(CO)n(PMe3)4-n (n ¼ 2, 3) [(Z5-C5H5)2Fe2(CO)3]a Fe(CO)4 Ru(CO)2(PMe3)2 Group 9 Tp’Rh(L) (Tp’ ¼ Tp , Tp4-tBu-3,5-Me; L ¼ CO, CNCH2CMe3) (Z5-C5R5)Rh(CO) (R ¼ H, Me) RhCl(PR3)2 (R ¼ Ph, p-Tol, Me)

Alkane

References

n-heptane

[21]

various methane cyclohexane n-pentane methane

[12,22] [23] [24] [25] [26]

various methane methane various various methane cyclopentane various

[27] [7a] [7a] [14,28] [29] [7a] [30] [31]

various various various methane

[32] [33] [22b,34] [35]

various cycloalkanes cyclohexane

[19,36] [37] [38]

a

This fragment was not shown conclusively to form an alkane complex. Bpin ¼ pinacol boryl; dfepe ¼ bis [bis(pentafluoroethyl)phosphino]ethane; dmpe ¼ bis[dimethylphosphino]ethane; Tp ¼ hydrotris(pyrazol-1-yl)borate; Tp ¼ hydrotris(3,5-dimethyl-pyrazol-1-yl)borate; Kp ¼ cyclopentadienyltris(diethylphosphito)cobaltate; Tp4-tBu-3,5Me ¼ hydrotris(4-tert-butyl-3,5-dimethyl-pyrazol-1-yl)borate.

Fig. 4 Direct observation of an alkane C–H bond coordinated to a metal was reported by Ball, who was able to characterize [CpRe(CO)2(cycloC5H10)] by NMR spectroscopy at below 193 K. NMR spectroscopic characterization allows direct observation of the bound alkane hydrogen (left) and quantitative measurement of the weakened of the bound C–H bond (right). Adapted with permission from Geftakis, S.; Ball, G. E. J. Am. Chem. Soc. 1998, 120, 9953. Copyright 1998 American Chemical Society.

methylene C–H binding sites are thermodynamically preferred to methyl sites in the [CpxRe(CO)2] (Cpx ¼ cyclopentadienyl based ligand) systems studies by NMR.40 In contrast to the thermodynamic selectivity observed in [CpxRe(CO)2(alkane)] systems, various alkane activation studies using transition metals have observed preferential or exclusive methyl C–H activation.8a TRIR studies have found that there is a much lower enthalpy of activation for terminal C–H positions as compared to secondary positions.37 However, NMR observations of alkane complexes have been able to further explain trends in C–H activation chemistry. For example, terminal binding of linear

Alkane s-Complexes

507

alkanes is initially dominant in [CpxRe(CO)2] type systems (i.e. they are the kinetic product), and it is observed that alkanes with more methylene groups give rise to larger concentrations of methylene bound isomers (based on the expected statistical distribution of products). These NMR observations support longer chain alkane s-complexes having longer average lifetimes prior to C–H activation in [CpxRh(CO)(alkane)] systems.37b Additionally, the bulky facial ligand hexaethylbenzene (HEB) results in a thermodynamic preference for terminal alkane binding in isoelectronic rhenium and tungsten systems {[(HEB)Re(CO)2]+ and [(HEB)W(CO)2]}.24b,41 Ball has suggested that this selectivity arises based on an endergonic penalty incurred with longer chain alkanes needing to adopt gauche configurations in the presence of the bulky HEB ligand. This corresponds to lower average lifetimes for long-chain linear alkane rhodium s-complexes bearing Cp ligands as compared to Cp ligands. Thus, NMR observations help aid our understanding of the relative rates of terminal C–H activation mediated by transition metal complexes. NMR studies on the family of rhenium alkane s-complexes of the type [LpRe(CO)2(alkane)] have been extended to include a variety of cyclopentadienyl derived ligands and Cp analogues trispyrazolylhydroborate (Tp) and cyclopentadienyltris(diethylphosphito)cobaltate (Kp) (Fig. 5).30,31 Additionally, Ball incorporated the neutral facial ligand HEB onto a rhenium dicarbonyl fragment to generate cationic alkane complexes of the type [(HEB)Re(CO)2(alkane)][Al(OC(CF3)3)4] (alkane ¼ cyclopentane, n-pentane).41 The stability of cationic alkane complexes was found to be superior to those of the previously reported neutral complexes. Besides the family of rhenium based alkane complexes, isoelectronic manganese and tungsten d6 complexes have been reported (Fig. 6).24b,29 Ball was able to generate tungsten complexes of the type [(HEB)W(CO)2(alkane)] (alkane ¼ n-pentane, cycloheptane, isobutane, 2,2-dimethylbutane).24b Although these were less stable than the rhenium family of alkane complexes, zerovalent group 6 complexes were the first complexes to be postulated as alkane complexes with IR evidence. Importantly, as spin-1/2 183W is 14.3% naturally abundant, 183W–1H and 13C–1H couplings could be observed in the same sample, confirming the validity of a three center two electron bonding model. Perutz and George were also able to obtain NMR spectroscopic data for the manganese alkane complexes [CpMn(CO)2(alkane)] (alkane ¼ ethane, propane, n-butane, cyclopentane, isopentane).29 As the stability of alkane complexes decreases from 5d to 3d metals, much lower temperatures were required to obtain NMR data. Due to solvent limitations, experiments were necessarily conducted in liquid ethane, propane or n-butane.

Fig. 5 Summary of rhenium s-complexes characterized by NMR spectroscopy.14,28b,30,31,39–42

Fig. 6 Summary of tungsten and manganese alkane s-complexes that are isoelectronic to [CpRe(CO)2(alkane)] type complexes that have been observed by NMR spectroscopy.24b,29

508

Alkane s-Complexes

Fig. 7 Rhodium methane and ethane complexes characterized by NMR spectroscopy. In contrast to the other alkane complexes observed by NMR, these complexes were synthesized by protonation of rhodium methyl and ethyl precursors (Fig. 2C) as opposed to a photolytic synthetic method (Fig. 2B).16

Apart from the characterization of photolytically generated d6 group 6 and group 7 alkane complexes, Brookhart has characterized rhodium methane and ethane complexes (Fig. 7) generated via protonation of alkyl ligands (Fig. 2C).16 Although the products were thermodynamically unstable, the alkane complexes (i.e. the kinetic products) were sufficiently long lived at low temperature to allow detailed NMR analyses. Brookhart also established simultaneous 103Rh–1H and 13C–1H couplings in [Rh(PONOP)(alkane)] [BArF4] systems. The isotope 103Rh is spin-1/2 and 100% naturally abundant, simplifying the identification of metal-ligand coupling. The barrier for the interchange of coordinated geminal alkane hydrogen atoms for these complexes was very small giving rise to time averaged signals, however, the barrier for metal migration between vicinal hydrogen atoms (i.e. 1,2-migration) was 7.2 kcal mol−1 at 161 K, allowing the observation of a fundamental exchange process termed ‘chain-walking’ where a bound alkane exchanges coordinating C–H positions between neighboring carbon positions. Weller has also obtained solid state NMR (SSNMR) data on crystalline samples of rhodium alkane complexes (discussed below).17b,c,f SSNMR provided direct detection of 31P and 13C NMR signals of the alkane complexes under study, however, SSNMR provides very low resolution 1H NMR data precluding direct detection of C–H


Rh interactions via 1H NMR spectroscopy, and preventing the observation of 1JCH data to provide insight to the strength of the metal-alkane interaction. However, indirect methods based on heteronuclear correlation spectroscopy (HETCOR SSNMR) allow chemical shifts for coordinated protons to be derived. The SSNMR techniques employed by Weller can be of most use to determine reaction yields, complex stability, and decomposition kinetic data, but provide less decisive evidence for alkane metal interactions, as these are better indicated by observed 1JCH rather than chemical shift. As such, solution NMR has a clear advantage in being able to provide information on the extent to which a C–H bond is ‘stretched’ through observation of bound C–H groups and their respective coupling constants.

1.16.2.2 1.16.2.2.1

Structurally characterized alkane s-complexes Single crystal X-ray diffraction (SCXRD)

The benchmark structural characterization method in chemistry is currently single crystal diffractometry. In most cases, a diffraction pattern is obtained using an X-ray source, although electron and neutron sources are also feasible. Single crystal X-ray diffractometry (SCXRD) offers the ability to generate high precision molecular structures for molecules that stack in a crystal lattice in the solid state. Indeed, much of our understanding of the electronic structure of molecules and bonding arises from insights into molecular geometry provided by SCXRD. The verification of s-complexes of hydrogen, borane, alane, silane and stannane was obtained crystallographically with SCXRD studies and/or neutron diffraction studies.6 In large part, the solid state characterization of non-alkane s-complexes was found to correlate well with solution spectroscopic studies. Despite the ever-growing spectroscopic evidence that helps us understand the solution behavior of alkane s-complexes, there exists no example of an alkane s-complex that has been characterized both the solid state and in solution. Indeed, the only solid state molecular structures of alkane s-complexes to be substantiated with other characterization techniques belong to the family of complexes synthesized by Weller, which have also been characterized by solid state NMR.17b,c,f This section summarizes reported structures that contain close metal-alkane contacts. In all cases, the ligand structure supporting the metal greatly assists alkane coordination. Nonetheless, a great deal about the nature of alkane binding to metals can be derived from the reported structures if the context of the ligand/anion/crystal lattice structure is taken into account. The first report of a close alkane-metal interaction was reported by Reed and Boyd in 1997 (Fig. 8).10a They reported the inclusion of a n-heptane molecule in a double A-frame iron porphyrin complex, [Fe(DAP)(n-C7H16)], generated from recrystallization of the porphyrin complex in fluorobenzene/n-heptane solution. The DAP ligand creates a cavity that supports the inclusion of the n-heptane molecule, nonetheless, short Fe–C distances of 2.52(4) Å and 2.83(3) Å are observed from the two disordered iron

Alkane s-Complexes

509

Fig. 8 Molecular structure of the iron n-heptane complex [Fe(DAP)(n-C7H16)] reported by Reed and Boyd.10a

sites to the closest n-heptane approach. In addition to the iron sites being disordered, the n-heptane is also disordered over two sites, preventing accurate location of any alkane hydrogen positions. The two reported Fe–C distances representing the closest n-heptane approaches to each iron site lie either side of the maximum reasonable covalent bond distance calculated for a linear Fe–H–C style interaction.43 Indeed, preliminary DFT calculations supported the binding of C1–C4 alkanes to the iron center with binding energies ranging from 10.5–16.7 kcal mol−1. At the time of the report, the only other method of alkane s-complex detection had been via FTIR and UV-Vis spectroscopies. Corroboration of the alkane ligand in solution would be difficult using such techniques, and although the presence of free n-heptane was established through NMR analysis, no attempt was made to observe an alkane-iron interaction via this method. In 2003, Meyer reported a series of uranium(III) complexes that incorporated an alkane molecule of recrystallization (Fig. 9).10b The uranium center was supported by the bulky tris-amino-tris-aryloxide ligand, (ArO)3tacn. The uranium complex [U((ArO)3tacn)] incorporates a single molecule of methyl-cyclohexane, methyl-cyclopentane, cyclohexane, cyclopentane or neohexane when crystallized in the presence of the appropriate alkane and n-pentane. The preferential binding of cycloalkanes and methylene groups would seem to agree with NMR observations reported for rhenium systems,40 however, a competition experiment between neohexane and cyclopentane resulted only in formation of the neopentane adduct, suggesting that binding may be attributed to the ability of the alkane to fit the lipophilic pocket size created by the (ArO)3tacn ligand. Indeed, it was observed that thermal noise for the alkane carbon atoms tended to be smallest for the carbon atoms closest to the uranium center. This can be interpreted as stabilization of these carbon positions through binding of the alkane to uranium center, or through ‘squeezing’ of the alkane in the tapered pocket size. The observed distances between the uranium center and the closest alkane carbon atom approaches ranges from 3.71(1)-4.04 (3) Å, distances that fall well outside the maximum reasonable covalent bond distance of a U-H-C linear interaction.43 However, these distances are well within the sum of the van der Waals radii of uranium and the closest hydrogen atoms,44 suggesting that an electrostatic interaction between the alkanes and uranium is feasible. Preliminary DFT studies suggested a binding energy of the alkanes to the uranium center in the order of 11 kcal mol−1, similar to those found for Reed’s Fe complex, however, analysis of the s-interaction found less than 2% contribution from alkane based s/p orbitals. Given the ‘hard’ nature of U(III) and the low C/H

Fig. 9 Molecular structure of the uranium methyl-cyclohexane complex [U((ArO)3tacn)(cycloC6H11(CH3))] reported by Meyer.10b

510

Alkane s-Complexes

Fig. 10 Molecular structure of the potassium n-hexane complex [K2(DAP)(n-C7H16)] reported by Emslie.10c

contribution to any s-interaction, it is highly likely that the uranium-alkane interactions can be considered a combination of electrostatic metal-alkane interactions supported by the lipophilic ligand pocket. A series of alkane ligands bound to potassium ions reported by Emslie in 2013 were also found to be largely electrostatic interactions (Fig. 10).10c Emslie found that incorporation of two potassium cations into a diaminoxanthene based pincer ligand (XAT) framework, and subsequent crystallization from toluene/alkane solvent mixtures led to structures of the type [K2(XAT) (alkane)], where an alkane was found in close proximity to one of the potassium centers. Similar to the previous structures mentioned, the XAT ligand provides a protective pocket that accommodates an alkane molecule. Although the alkane potassium closest approach K–C distances of 3.22(1)-3.48(3) Å are mostly below the expected covalent bond distance for a linear K–H–C interaction,43 the interaction was found to be predominantly electrostatic in nature when modelled with DFT. In 2012 Weller reported the first of the most comprehensively reported classes of alkane s-complexes (Fig. 11).17,18 Rather than relying upon serendipitous incorporation of an alkane molecule through crystallization approaches, Weller actively generates the alkane ligand in the metal’s coordination sphere through the hydrogenation of alkene ligands (Fig. 2 method D). Additionally, Weller has been able to reliably generate X-ray quality crystals using crystal-to-crystal solid-state (SS) synthetic methods. These have been largely successful due to the similarity in unit cells of the starting material metal-alkene complexes and alkane s-complex products, allowing preservation of the samples’ crystallinity. Weller has (thus far) only applied this method to group 9 cationic complexes (rhodium and cobalt), however, the method may in principle be extended to other cationic metal alkene complexes. The importance of the counteranion used by Weller {[BArF4]−, tetrakis(3,5-bis(trifluoromethyl))phenylborate} has been noted due to its ability to form a supporting octahedral (Oh) lattice that retains the alkane ligand, while admitting hydrogen for the purpose of alkene hydrogenation.17b Notably, the anion structure allows free movement of the alkane within the individual anion protected cavities, providing a range of coordination modes of starting materials, intermediates, and alkane complex products. This differs greatly from previously discussed structures, where the binding mode of the alkane is greatly influenced by the shape of the metal’s supporting ligand pocket that supports the alkane ligand. Weller’s original report dealt with the coordination of norbornane (NBA) to a cationic bis-phosphinorhodium(I) fragment (Fig. 11).17a Indeed, the majority of Weller’s reports have been concerning s-NBA rhodium(I) complexes.17a,b,d–j Weller has also reported coordination of tert-butane, n-pentane and cyclohexane to rhodium(I), cyclooctane and NBA to rhodium(III), and NBA to cobalt(I).17b,c,18 The alkanes that Weller has employed allow the distinct advantage of binding the metal in a chelate (bidentate) fashion, providing additional thermodynamic stability (Fig. 11). In contrast to other isolated reports of alkane s-complexes, the metal-carbon distances in Weller’s structures are remarkably short, and well within a reasonable covalent bond distance.43 Indeed, the range of Rh–C distances for Weller’s alkane s-rhodium(I) complexes is from 2.36(1)-2.53(6) Å, while the precursor rhodium(I) alkene complexes display only slightly shorter Rh–C distances in the range 2.136(9)-2.329(9) Å, and an agostic interaction in a rhodium isobutene precursor is observed at 2.368(9) Å. In sharp contrast to other examples of solid state structures, DFT analysis of Weller’s rhodium(I) alkane s-complexes confirms considerable metal-alkane binding energies on the order of 13.5–18.5 kcal mol−1.17b–d Although these binding energies are notably less than calculated alkane binding energies in the solution stable rhenium alkane s-complexes discussed above (20–25 kcal mol−1),41 DFT has also correlated additional support from the anion and crystal lattice environment through dispersive stabilization.

1.16.2.2.2

Neutron diffraction studies

Historically, neutron diffraction has served as a complementary technique to X-ray diffraction, where neutron sources are much more scattered by lighter elements than X-rays. Thus, many structures that required accurate location of hydrogen atom positions employed neutron diffraction to confirm the absolute structures. Indeed, the seminal dihydrogen s-complexes synthesized by Kubas were confirmed using a combination of X-ray and neutron diffraction analysis.4a Methods in SCXRD have evolved to be able to much better locate light atoms such as hydrogen, even in close proximity to heavy atoms. Nonetheless, neutron diffraction not only offers superior scattering by light atoms, but is able to distinguish between 1H and 2H positions well due to the large difference

Alkane s-Complexes

511

Fig. 11 Molecular structure of the rhodium norbornane (NBA) complex [Rh(Cy2P(CH2)3PCy2)(NBA)][BArF4], and a summary of all the group 9 alkane complexes reported by Weller.17,18

in the incoherent scattering of hydrogen and deuterium. Due to this, where possible, many neutron diffraction experiments are run using deuterated samples to avoid the large incoherent scattering arising from hydrogen. The synthetic method of generating alkane in a metal coordination sphere employed by Weller lends itself to the late-stage incorporation of deuterium into precursor olefin ligands, and active C–H sites. Weller has been able to demonstrate the high dynamic flexibility of NBA within the anion ‘nanoreactor sites’ through both treatment of precursor norbornadiene complexes with D2 (as opposed to H2), and also by the treatment of the already synthesized alkane s-complex [Rh(Cy2P(CH2)2PCy2)(NBA)][BArF4] with D2.17g As the alkane complex was able to reversibly eliminate H2, D2 incorporation into the NBA ligand was possible. Weller used a combination of neutron diffraction of the deuterated alkane complexes and 2H NMR analysis of the liberated norbornane after the alkane complex had been dissolved to identify which positions had incorporated deuterium. The positions of the deuterated sites demonstrated that the alkane ligand was able to ‘pivot,’ ‘rotate,’ ‘rock’ and ‘tumble’ within the anion protected cavity. In 2012 Long reported the structures of Fe-MOF-74 {Fe2(dobdc), (dobdc)4− ¼ 2,5-dioxido-1,4-benzenedicarboxylate} with ethane-d6 and propane-d8 adsorbed onto the MOF’s internal surface as shown by neutron powder diffraction techniques (Fig. 12).45 Similar magnesium, copper and chromium based MOFs have also been reported that exhibited methane adsorption.46 However,

512

Alkane s-Complexes

Fig. 12 Molecular structures of the iron ethane and propane MOF structures [Fe2(dobdc)(C2H6)] and [Fe2(dobdc)(C3H8)] reported by Long.45

the adsorption of methane in these systems seems to be a function of the topology of the MOF rather than an alkane-metal covalent interaction, and binding in vacant ‘window’ cavities in the case of Cu/Cr was preferred over binding in vacant metal sites.46b Although powder diffraction solutions are becoming routine for structural characterization,47 this method is still considered to be less accurate than single crystal diffraction methods. Nonetheless, Long was able to experimentally determine the positions of the deuterium and carbon atoms in close proximity to five coordinate iron centers. The weak interaction between the deuterated light alkanes and the iron center is evident with elongated Fe–C distances (ethane ¼ 3.09(2) Å; propane ¼ 2.90(3) Å; maximum covalent distance ¼ 2.70 Å). The alkanes likely bind through an ‘ion induced dipole interaction,’ similar to that observed in methane adsorption onto an analogous magnesium MOF.46

1.16.2.3

Alkane complexes studied computationally

Alkane s-complex computational analyses can be classified as either studies that enhance our understanding of observable alkane complexes (i.e. complexes described above), or they can be studies of alkane complexes in systems that have no reported experimental data. Given that almost all complexes that undergo C–H activation are thought to be preceded by alkane complex intermediates,48 systems that study alkane complexes as intermediates enroute to C–H activation without experimental observation of such alkane complexes will be ignored here. Computational chemistry has found many important roles in the field of alkane s-complexes (Fig. 13). Indeed, DFT allows comparison between solution and solid-state alkane s-complexes. Despite the utility of computational chemistry in studying

Fig. 13 Computational chemistry plays an important role in our understanding of alkane s-complexes. (A) DFT can help refine/resolve/identify experimental data, and experimental data in turn helps to calibrate DFT models. (B) Calibrated DFT models can provide further predicted properties of an alkane complex that are not observable, or that have not been observed. (C) The nature of the alkane-metal bond can be understood with the aid of DFT (charge populations, bond critical points, atomic orbital contribution, etc.). (D) As yet unreported alkane complexes can be studied and their stabilities assessed in comparison to known systems. This can lead to experimental chemists concentrating their efforts on such complexes identified through DFT prediction.

Alkane s-Complexes

513

alkane complexes, there exist few stand-alone computational studies of alkane complexes.49 However, many experimental results have been supplemented with DFT analyses to verify and further understand the experimentally observed data. DFT studies have been able to help identify and resolve experimental data from alkane s-complexes (Fig. 13A). For example, poorly resolved 1H SSNMR signals arising from rhodium NBA complexes were identified by Weller and MacGregor and their cross-peaks resolved in correlation spectroscopy through the assistance of computational modeling.17a–d,f–i DFT has also helped resolve Rh–H bonding distances in solid state structures reported by Weller; while low electron densities at H prevent accurate determination of hydrogen locations using X-ray crystallography, DFT can provide more precise structural models.17c,d The models can also be calibrated to experimentally observed heavy-atom positions, or spectroscopic data. This can be extended to the prediction of experimental properties that haven’t been observed (Fig. 13B). Most notably, structural predictions can be made based on calibration to data that would not traditionally provide stereochemical information. For example, combined TRIR and DFT studies have been able to develop model simulations that not only can recreate the TRIR data, but provide structural information. In this way, George and Hall were able to interpret C–H activation parameters derived from TRIR data to understand periodic trends.37b Their simulated model suggested that these trends arose from the stability of intermediate alkane complexes and their ability to undergo 1,2-, 1,3- or 1,4-shifts to isomerize between different alkane coordination positions, and to avoid gauche configurations in doing so. The simulated model was able to emulate the experimentally observed activation parameters, providing further confidence in the structural information it provided. DFT predictions can be extended to fundamental thermodynamic parameters, such as bond energies and barriers for transformation between isomers, but they can also be applied to theoretical concepts that cannot be measured directly (Fig. 13C). For example, parameters such as bond critical points (BCP), natural bond orbitals (NBO), noncovalent interactions (NCI), charge densities and frontier orbital isosurfaces can be calculated using various DFT methods such as quantum theory of atoms in molecules (QTAIM) or periodic DFT, and provide a better understanding of the nature of the alkane-metal interaction. Importantly, the prediction of parameters, such as bond energies, based solely on the metal-alkane interaction can be compared to experimentally obtained values to ascertain the degree of stabilization that may be offered from additional factors such as the encompassing ligand or crystal environment.17a–d,f–i To this end, DFT has allowed the verification of metal-alkane bonding as opposed to alkanes simply being in close proximity to metal centers and stabilized through other non-metal based interactions. Lastly, DFT has proven powerful at predicting the stability of experimentally unreported alkane s-complex systems, thus guiding experimental chemists’ efforts in applying particular metal fragments in alkane coordination chemistry (Fig. 13D). Simulation of unknown systems can be attractive to efficiently screen untried metal fragments/ligand environments in a cost effective manner, given investment that must be made (in terms of time and resources) to experimentally study any alkane system. Indeed, computation studies and predictions of increased stabilities for both [TpRe(CO)2(alkane)] and cationic rhenium alkane systems led to experimental reports.30,41,49b,50 DFT has also allowed difficult-to-isolate alkane complexes to be studied in silico, where short lifetimes, side reactions, and technical difficulties associated with the synthesis and characterization of alkane complexes can be avoided. In this manner, many s-methane complexes have been studied through DFT despite the difficulties in working with methane gas in an experimental setting.49a,e,50

1.16.3

Bonding trends in alkane s-complexes

1.16.3.1

Bonding modes

The ability of alkanes to act as ligands has been reported since 1962, but it is only recently that experimental data have been available to reveal the way in which alkanes bind metals (i.e. stereochemical and structural information). Initial reports using fast-spectroscopic techniques offered no insight as to how the alkane interacted with reactive metal centers, and were at the time assigned as solvation effects.12a Structural models for how C–H bonds interact with transition metals were constructed upon knowledge obtained from related C–H agostic interactions and other isolable s-complexes.1 However, in the case of agostic interactions, rigidity arising from the chelate ligand, and in the case of other s-complexes, electronic differences in the E-H bond both provide limitations in their ability to model the way in which free alkanes might bind transition metal centers. Prior to any NMR or molecular structure reports, Hall and Perutz proposed a number of possible alkane binding modes (Fig. 14), which have come to define the nomenclature of this field.7a With the advent of NMR and SCXRD characterization, a great deal of knowledge has been furnished regarding alkane-metal bonding in solution and in less restricted solid-state environments, thus allowing improved understanding of alkane binding. For example, Modes 1–3 tend to be favored in electrostatic dominated interactions, while Mode 4 provides maximum overlap between alkane based s molecular orbitals and s antibonding orbitals with metal based d-orbitals (Fig. 14). NMR and SCXRD cannot only reveal the mode of alkane coordination, but also the degree of coordinated C–H bond weakening/elongation, allowing correlation between bonding modes, binding strength and the relative degree of C–H activation (i.e. how far a C–H bond is along the reaction pathway towards activation). As described above, alkanes can bind to transition metal centers in a similar manner to the Dewar-Chatt-Duncanson model outlined for metal-alkene bonding.3 This bonding description is composed of two distinct metal-alkane interactions; (i) s-donation of alkane electrons to a vacant metal orbital, and (ii) p-retrodonation of metal based electrons to an alkane s antibonding orbital. The preferred geometry of these two interactions does not always coincide, and thus the degree of one interaction as compared to the other can dictate the alkane binding position. For example, Mode 1 induces very little orbital overlap between empty metal

514

Alkane s-Complexes

Fig. 14 Possible alkane coordination geometries and nomenclatures, as first described by Perutz.7a

orbitals and C–H s orbitals so would be expected to have less stability than the other modes, however, it is observed in Weller’s rhodium(III) and Emslie’s potassium alkane complex structures where electrostatic interactions may dominate.10c,17d Somewhat unsurprising, solution studies have found that the stability of alkane complexes greatly increases with the period and size of transition metal centers as such metals allow stronger p-retrodonation.21c This explains the relative stability of rhenium alkane complexes, with [CpRe(CO)2(n-C7H16)] being ca 1000 times more stable than [W(CO)5(n-C7H16)], ca 300 times more stable than [CpMn(CO)2(n-C7H16)] and astoundingly ca 50,000 times more stable than [CpV(CO)3(n-C7H16)].

1.16.3.2

Bonding distance and angles

Fig. 15 illustrates the reported bonding distances in solid-state molecular structures, and that obtained via XAFS by George. The dotted line on the graph represents conservative estimates for the maximum covalent bonding distance (i.e. the assumption of a linear M–H–C covalent interaction) between a C–H bond and a metal center. It is evident that many of the reported interactions fall near or outside of this limit. Indeed, only Weller’s rhodium(I) and high-spin cobalt(I) alkane complexes have high confidence level bonding distances comfortably below this limit for covalent bonding.17,18As such, it is evident that further analysis would be warranted in cases with low precision to provide greater resolution of actual metal-alkane bond distances and angles. Nonetheless, this is not to say that significant interaction between alkane and metal is not present in all of the discussed examples. Indeed, the van

Fig. 15 Minimum carbon-metal distances for alkane ligands from all of the structurally reported examples described above. The diagonal represents the calculated maximum covalent distance for a linear M–H–C interaction. Maximum covalent distance calculated from the sum of the metal covalent radius, an sp3 carbon covalent radius and two hydrogen covalent radii (i.e. maximum covalent distance assumes a linear interaction). Covalent radii of metal centers, carbon and hydrogen taken from reference [43a].

Alkane s-Complexes

Table 2

515

Summary of bonding distances and angles from structurally characterized alkane complexes.

Alkane complex set

M-C distance range (Å)

M-H-C angle range ( )

References

[Fe(DAP)(n-heptane)] (Reed) U(III) alkane complexes (Meyer) Rh(I) alkane complexes (Weller) Rh(III) alkane complexes (Weller) Co(I) NBA complex (Weller) K(I) alkane complexes (Emslie) Fe(II) alkane complexes (Long) [W(CO)5(n-heptane)] (George)

2.52(4)-2.83(3) (2.68) 3.71(1)-4.04(3) 2.36(1)-2.53(6) 2.75–2.90(9) 2.61(1) 3.22(1)-3.48(3) 2.90(9)-3.09(6) 3.07(6)

N/A 100.6–128.5 (119.0) 98.0–131.8 (109.8) 141.0–146.0 (143.5) 139.9 101.5–118.2 (112.4) 107.7–123.5 (115.6) N/A

[10a] [10b] [17] [17d] [18] [10c] [45] [20]

der Waals radii limit (4.67–5.00 Å) in each of the reported cases well exceed the observed bonding distances taking into account reported metric standard deviations.44 Table 2 summarizes the reported distances for each of the alkane complex systems reported, along with the reported range of angles and the average of the alkane complex set. The high degree of freedom of alkane ligands in Weller’s rhodium alkane complexes allows comparison of the metal-alkane bonding character through the observed Rh–H–C angles exhibited by the coordinated C–H bonds. For a rhodium-Z2-C,H covalent interaction (Mode 4, Fig. 14), an angle that approaches 90 is preferred to allow maximum orbital overlap, as established by the Dewar-Chatt-Duncanson model. An electrostatic (Z1-H) interaction will prefer angles approaching 180 (Mode 1, Fig. 14). It can be seen from the majority of Weller’s rhodium(I) complexes that a more acute angle is preferred (Rh–H–C mean ¼ 109.8 ) as compared to the rhodium(III) alkane complexes he has reported that prefer a more obtuse angle and an Z1-H coordination mode (Rh–H–C mean ¼ 143.5 ). Due to the more rigid restraints of the ligand environment supporting the coordinating alkane in all other reported structures, it is difficult to judge the strength of orbital overlap in the metal-alkane interaction through the M–H–C angle. However, it was concluded that electrostatic interactions dominate in Emslie’s potassium-alkane adducts,10c and in Long’s and Yildirim’s examples of alkanes adsorbed on the interior surface of iron and magnesium MOF structures.45,46 This is somewhat supported by the C–H metal bonding distances being near or greater than the limit of covalent bonding (Fig. 15). Bonding interactions between metal centers and alkanes in solution are estimated to be generally larger than those found in solid state structures. Although the tungsten-carbon distance obtained by George using XAFS lies slightly outside the calculated covalent bonding limit, the degree of error using this technique is quite high, and the distance reported represents all the different structural isomers present in solution.20 Comparison to DFT predictions for the average n-heptane binding distance to the ‘W(CO)5’ fragment for various isomers were in strong agreement with the XAFS result. As such the minimum bonding distance between n-heptane and tungsten may be closer to the DFT predicted distance of 2.86 Å for methyl site binding, and within the maximum covalent limit for a C–H–W interaction (viz. 3.00 Å).

1.16.3.3

Alkane complexes versus fluxional alkyl hydrides

The observation of alkane s-complexes prior to C–H activation is reported,19 as are examples of alkane s-complexes in equilibrium with their C–H activation products.28b As such, alkane s-complexes have been validated as intermediates en route to transition metal mediated C–H activation. However, many observable alkane s-complexes, such as the complexes [CpxRe(CO)2(alkane)], do not lead to C–H activation. Indeed, this has allowed their in-depth study using slower spectroscopic measurements. The ability to differentiate stable alkane complexes versus alkane complexes in equilibrium with their C–H activation products has been made possible through a combination of the characterization methods described above. Weller used simple deuteration studies to demonstrate that NBA in his [Rh(PCy2(CH2)2PCy2)(NBA)][BArF4] system was undergoing unobserved C–H activation.17g However, such an approach only confirms the existence of C–H activation and alkane complex intermediates, and neither reveals kinetic or thermodynamic data regarding the two isomeric forms. Alkane complexes bearing carbonyl ligands have proven very sensitive to changes in the formal oxidation state of the metal upon C–H activation. Indeed, Bergman and Harris relied upon such data to assign alkane complex intermediates and C–H products in the study of alkane C–H by a ‘Tp Rh(CO)’ fragment.19 Obviously, other processes are possible that affect the metal electron density and consequently the carbonyl stretching frequency, and IR spectroscopy cannot monitor the alkane C–H bonds directly, as such IR methods are not conclusive in determining, or differentiating between, C–H activation products and alkane complexes. In the event of a fast exchange between free alkane, an alkyl hydride complex and an alkane complex, NMR may observe a timeaveraged signal that results in signals that are falsely indicative of an alkane complex.28b For example, metal bound C–H bonds give rise to 1H NMR resonances in the range −2 to −6 ppm. However, time-averaged signals arising from free alkane and alkyl hydride exchange would also be expected to give rise to 1H NMR signals in this region. Two methods have been proposed to distinguish between alkane complexes and alkyl hydrides in fast exchange with free alkanes by NMR spectroscopy. The first method relies upon the relatively large C–H coupling constant (1JCH) of a bound C–H bond as compared to the much smaller C–H coupling constant of an alkyl hydride (3JCH). In the case of an alkane complex, 1JCH for the

516

Alkane s-Complexes

bound C–H bond would be expected to above 90 Hz, whereas 3JCH in metal alkyl hydride complexes is generally below 20 Hz. Thus, the time-averaged 1H NMR signals of bound C–H bonds in alkane complexes tend to have very large JCH.11 The second method relies upon the observation that alkane complexes tend to have very large Isotopic Perturbation of Resonance (IPR) effects, and that C–H bonds are greatly preferred in metal coordination over C–D bonds. IPR effects arise from the zero point energy difference between terminal C–H and C–D vibrational modes, which is larger than that between the bound C–H–M and C–D–M interactions, and therefore contributes more to the sum of the bond energies. This is despite the fact that the C–D–M interaction is expected to be stronger as compared to that of C–H–M. Thus C–D bonds have a thermodynamic preference to occupy terminal positions over C–H bonds, giving rise to a preference for C–H coordination.40,42 In simplistic terms, this can be thought of as placing the C–D bond in a tighter potential energy well (terminal position), and the C–H bond in a flatter potential energy well (bridging position) as is illustrated in Fig. 16. Large IPR effects in partially deuterated alkanes with geminal protio/deuterio positions are evidence for a Z2-C,H binding mode (Mode 4, Fig. 14) not possible in alkyl hydride complexes given that this effect is contingent upon differing zero point energies of terminal and bridging C–H/D bonds.30,40,51 Indeed, given that such an effect would be absent in methylene groups (H2CR2) binding via mode 2 and methyl groups (H3CR) that bind via mode 3 (as these bonding scenarios would be devoid of terminal C–H/D bonds), IPR effects observed in such systems provides evidence for bonding through Modes 1 or 4 (Fig. 14).

1.16.3.4

Binding selectivity and dynamic exchange processes

Transition metal mediated alkane C–H activation shows a clear preference for methyl C–H bonds over methylene and methine C–H bonds, or in the words of Labinger and Bercaw, “Transition-metal centers preferentially activate terminal C–H bonds”.8a Indeed, TRIR studies have shown that activation enthalpies for primary C–H bonds are much lower than those of secondary psotions.37b Previously, although it was clear that methyl positions were preferentially activated, it was not understood or observed whether methyl bound or methylene bound alkane s-complexes were thermodynamically preferred, given that fast spectroscopic techniques could not discern between the isomers. However, the advent of NMR and solid state structural studies allow us to examine the stereochemistry of alkane s-complexes prior to activation, and in particular discriminate between different binding sites in the same alkane ligand. Such observations can help explain the longer lifetimes of longer chain alkane s-complexes prior to activation, in that methylene bound isomers with higher bond activation enthalpies are present in higher concentrations. Apart from determining the thermodynamic preference of different alkane binding sites, NMR allows the extraction of kinetic data for processes that account for binding site exchange. Brookhart has calculated exchange rates between adjacent methyl positions in a rhodium ethane s-complex using VT 13C NMR line shape analysis (Fig. 17A).16b The observed 1,2-migration process has been labelled ‘chain-walking,’ and interchange between alkane binding sites is often attributed to this process. Ball has used 2D ROESY NMR to observe 1,2-shifts in [CpiPrRe(CO)2(n-C5H12)] (CpiPr ¼ isopropylcyclopentadienyl) methyl and methylene bound isomers.40 He also used this approach to observe exchange between bound methyl positions in a tungsten n-pentane s-complex (Fig. 17B).24b In this instance, no methylene bound isomers were observed, and Ball suggested that the methylene isomers eluded NMR observation due to fast ‘chain-walking.’ However, calculations by Hall and George have suggested that the barrier to 1,3- and 1,4-migrations can be similar to or smaller than 1,2-migration processes, so methylene isomers may not be mandatory intermediates in the exchange of two alkane terminal positions.37b Indeed, Ball later reported the observation of both 1,2 and 1,3-migration in the cyclopentane ligand of [Re(HEB)(CO)2(cycloC5H10)][Al(OC(CF3)3)4] (Fig. 17C).41 In this instance, cross-peaks in ROESY and NOESY NMR experiments were not well resolved and Ball instead used 1D EXSY NMR experiments to provide kinetic data.

Fig. 16 Isotopic perturbation of resonance (IPR) effects are observed in alkane complexes with equivalent hydrogen (C–H) and deuterium (C–D) sites. The effect arises from the ability of a terminal C–D bond to achieve a lower zero point energy (ZPE) due to its lower vibration frequency. Thus, the overall energy of a terminal C–D bond and a bridging C–H–M bond is lower than the energy of a terminal C–H bond and a bridging C–D–M bond.

Alkane s-Complexes

(A)

517

(B)

(C)

Fig. 17 (A) 1,2-migration between vicinal methyl positions in a rhodium ethane complex.16b (B) Exchange of terminal n-pentane positions in a tungsten n-pentane complex.24b (C). 1,2- and 1,3-migration observed in a rhenium cyclopentane complex.41 Note: Counter anions omitted in (A) and (C).

The solid-state molecular structures discussed above show a variety of binding selectivity dependent upon individual circumstances. Meyer’s series of uranium alkane structures were formed in presence of n-pentane solvent, but an n-pentane adduct was not observed.10b Methyl cyclohexane and methyl cyclopentane structures were shown to bind through methylene C–H bonds within the ring preferentially to methyl positions. However, a competition experiment between cyclopentane and neohexane established that neohexane, binding through a methyl position was the dominant alkane complex observed. Presumably, large lipophilic and host-guest interactions may well have attributed greatly to the preferred coordinating alkane in these structures. Reed’s iron n-heptane complex binds n-heptane through the methyl site, however, further penetration into the supporting A-frame porphyrin to bind a methylene site would result in either large steric repulsion, or an unfavorable twisting of the n-heptane molecule.1a Long’s report of iron ethane and propane adducts does show a closest approach of the propane molecule through the methylene group, however, this would also give rise to the largest surface interaction between the alkane molecule and the MOF support, thus the binding selectivity observed in this instance is unlikely to be due to preferred covalent bonding between methylene C–H bonds and iron (as compared to methyl C–H bonds and iron).45 Emslie’s report on a series of potassium alkane adducts shows binding through methyl groups, with the exception of a cyclopentane adduct that binds through a methylene, but was grown from cyclopentane solution (i.e. in the absence of methyl groups).10c The binding mode of these potassium alkane adducts not only reflects the preferable end on binding resulting from the electrostatic interaction, but is also supported by the cavity in which the potassium ion resides. Weller’s series of reports on group 9 alkane complexes, which offer the most convincing examples of covalent alkane-metal interactions, contain coordinated alkane molecules with methyl, methylene and methine groups.17,18 Although in many cases, the binding selectivity may be resultant on reducing bond strain in facilitating the alkane to bind in a chelate fashion, there are examples where certain site selectivity is evident. For example, n-pentane binds through methylene groups in the 2,4-positions, as opposed to binding through the 1,3-positions, which would provide similar bond strain in the alkane. Weller also observed that isobutane binds through methyl and methine groups, despite these positions being less geometrically aligned with the idealized observed square-planar geometry around the rhodium center as compared to binding solely through isobutane methyl groups. Based upon the very limited number of examples that show binding group preference independent of other factors (such as bond strain in forming square-planar chelate complexes), binding selectivity in Weller’s Rh(I) alkane complexes seems to suggest the methine and methylene C–H bonds are preferred over methyl C–H bonds. NMR studies have also allowed direct detection and elucidation of thermodynamic data for different alkane binding sites. Such data are most readily available for the widely studied neutral rhenium alkane complex family, [LpRe(CO)2(alkane)].14,28b,30,31,39,40,51 No data are available for Brookhart’s rhodium alkane complexes as methane and ethane alkane s-complexes both only contain one type of binding site.16 Binding preference in George and Perutz’s manganese alkane complexes for methyl and methylene positions could not be distinguished. In studies with propane, n-butane and isopentane, methyl C–H bound isomers are always observed as the major kinetic products, but the short lifetimes of both the methyl bound CpMn(CO)2(Z2-C1-H-alkane) and methylene bound CpMn(CO)2(Z2-C2-H-alkane) isomers prevented accurate assessment of equilibrium constants between these isomers.29 It must be noted that CpMn(CO)2(cyclopentane) was formed from reaction of cyclopentane with CpMn(CO)2(C3H8), however, this follows a general trend of heavier alkanes being more stable ligands than lighter alkanes and may not be indicative of binding site preference.21c,52 Interestingly, when employing isopentane, no methine bound isomer was observed (i.e. coordinated C2 position), however, the short lifetimes of these complexes may only reveal the kinetic reaction products.

518

Alkane s-Complexes

(A)

(B)

(C)

(D)

(E)

Fig. 18 Thermodynamic preferences for alkane binding with the ‘(Z5-C5H5)Re(CO)2’ fragment. Binding preference follows: (A) C–H > C-D40 (B) n-alkanes>cycloalkanes7c (C) CH2 > CH340 (D) C–Haxial > C–Hequatorial (for cyclohexane)51 (E) larger cycloalkane rings >smaller cycloalkane rings.7c These preferences are consistently observed for all other reported neutral rhenium alkane complexes.

Lastly, Ball has reported that tungsten alkane complexes have a large thermodynamic preference for methyl site binding over methylene sites. For example, his report of [W(HEB)(CO)2(n-C5H12)] notes that methyl binding is exclusively observed, although fast ‘chain-walking’ through methylene positions is highly likely as metal bound and free methyl sites of the n-pentane ligands were found to exchange.24b This observation may be specific to the HEB ligand environment rather than specific to tungsten, as an isoelectronic rhenium complex [Re(HEB)(CO)2(n-C5H12)][Al(OC(CF3)3)4] was found to prefer methyl binding,41 and Ball has shown in silico that methylene coordination of n-pentane requires the n-pentane ligand to adopt a less stable gauche configuration in the HEB ligand environment.24b Despite the above anecdotal reports, the bonding trends in neutral rhenium alkane complexes are consistent with a preference for methylene binding over methyl binding. Ball has also explored preferred binding between other inequivalent sites and between various alkanes. For example, a study of [CpRe(CO)2(cycloC6H12)] revealed that binding of axial C–H positions is preferred over equatorial C–H positions in geminal methylene sites.51 This observation strongly supports either Z1-H or Z2-C,H coordination modes (Modes 1 and 4, Fig. 14) in rhenium alkane complexes. Further, Ball has shown that cycloalkanes are preferred ligands to linear alkanes of similar molecular weights, and that larger cycloalkanes are preferred over smaller cycloalkanes, in agreement with previous IR based reports.7c,22e,52 Fig. 18 summarizes a number of dynamic processes and the preferred binding mode of the archetypal alkane complexes [CpRe(CO)2(alkane)].

1.16.4

Summary

As the sensitivity of analytical methods and the ingenuity of organometallic synthetic methods have improved, the observation and isolation of unstable complexes have become ever more possible. This is epitomized in the developing field of alkane s-complexes. Once considered highly unstable intermediates en route to C–H activation that were only observable through very fast timeresolved spectroscopic methods or theoretical modeling, alkane s-complexes are now routinely studied with IR, UV–Vis, XAFS and NMR spectroscopies in solution, and with X-ray and neutron diffraction and solid-state-NMR spectroscopy in the solid-state. Improved understanding of how alkanes bind to metal centers has allowed the evolution towards higher stability alkane complexes, which has further allowed in-depth studies to improve our knowledge of these complexes. Apart from the importance of studying alkane complexes as intermediates in processes such as alkane C–H activation or s-complex assisted metathesis (s-CAM), the field of alkane complexes is being driven from an intrinsic desire chemists have to expand our fundamental knowledge of weak metal-ligand bonding. The challenge of synthesizing an alkane complex that can be fully characterized in solution and the solid state still exists. In striving towards this goal, a host of new synthetic techniques and

Alkane s-Complexes

519

analytical approaches will continue to be developed that will benefit other related fields of organometallic chemistry. No doubt solution and solid state characterization approaches will be unified in a single complex at some stage in the near future, but for the field of alkane s-complex chemistry to continue to enjoy vigorous study, chemists in this field need to turn their attention to unique applications and novel chemistry that alkane s-complexes might allow. Certainly there are few better leaving groups than an alkane ligand, and after dissociation, alkanes are remarkably unreactive with other molecules. To this end, alkane s-complexes may be the best candidate chemists can devise for sources of metals with ‘vacant-sites.’ Thus, they may not only serve as catalytic intermediates in C–H activation processes, but as ideal pre-catalysts for a host of chemical processes.

References 1. (a) Kubas, G. J. In Comprehensive Organometallic Chemistry III; Mingos, M. P., Crabtree, R. H., Eds.; Elsevier, 2007; vol. 1; p 671; (b) Kubas, G.; Metal, J. Dihydrogen and s–Bond Complexes; New York: Kluwer Academic/Plenum Publishers, 2001. 2. Mingos, M. P. In Comprehensive Organometallic Chemistry; WiIkinson, G., Stone, F. G. A., Abel, E. W., Eds.; Pergamon Press: Oxford, 1982; vol. 3; p 1. 3. (a) Nelson, J. H.; Wheelock, K. S.; Jonassen, H. B. Chem. Commun. 1969, 1019; (b) Mingos, M. P. J. Organomet. Chem. 2001, 635, 1. 4. (a) Kubas, G. J.; Ryan, R. R.; Swanson, B. I.; Vergamini, P. J.; Wasserman, H. J. J. Am. Chem. Soc. 1984, 106, 451; (b) Crabtree, R. H. Chem. Rev. 2016, 116, 8750; (c) Kubas, G. J. Acc. Chem. Res. 1988, 21, 120. 5. (a) Norton, J. R.; Sowa, J. Chem. Rev. 2016, 116, 8315; (b) Connelly Robinson, S. J.; Heinekey, D. M. Chem. Commun. 2017, 53, 669; (c) Morris, R. H. Hydride Complexes of the Transition Metals in Encyclopedia of Inorganic and Bioinorganic Chemistry; John Wiley and Sons, 2011. 6. For seminal examples, see: (a) Hartwig, J. F.; Muhoro, C. N.; He, X.; Eisenstein, O.; Bosque, R.; Maseras, F. J. Am. Chem. Soc. 1996, 118, 10936;(b) Porschke, K. R.; Kleimann, W.; Tsay, Y.-H.; Kruger, C.; Wilke, G. Chem. Ber. 1990, 123, 1267; (c) Riddlestone, I. M.; Edmonds, S.; Kaufman, P. A.; Urbano, J.; Bates, J. I.; Kelly, M. J.; Thompson, A. L.; Taylor, R.; Aldridge, S. J. Am. Chem. Soc. 2012, 134, 2551; (d) Jetz, W.; Graham, W. A. G. Inorg. Chem. 1971, 10, 4; (e) Carre, F.; Colomer, E.; Corriu, R. J. P.; Vioux, A. Organometallics 1984, 3, 1272; (f ) Vincent, J. L.; Luo, S.; Scott, B. L.; Butcher, R.; Unkefer, C. J.; Burns, C. J.; Kubas, G. J.; Lledos, A.; Maseras, F.; Tomas, J. Organometallics 2003, 22, 5307; (g) Schubert, U.; Kunz, E.; Harkers, B.; Willnecker, J.; Meyer, J. J. Am. Chem. Soc. 1989, 111, 2572; (h) Piana, H.; Kirchgaessner, U.; Schubert, U. Chem. Ber. 1991, 124, 743. for reviews see:; (i) Lin, Z. Transition Metal s-Borane Complexes. In Contemporary Metal Boron Chemistry I. Structure and Bonding; Marder, T. B., Lin, Z., Eds.; Springer: Berlin, Heidelberg, 2007; vol. 130; (j) Lin, Z. Chem. Soc. Rev. 2002, 31, 239; (k) Schubert, U. Adv. Organomet. Chem. 1990, 30, 151 7. (a) Hall, C.; Perutz, R. N. Chem. Rev. 1996, 96, 3125; (b) Cowan, A. J.; George, M. W. Coord. Chem. Rev. 2008, 252, 2504; (c) Young, R. D. Chem. Eur. J. 2014, 20, 12704; (d) Weller, A. S.; Chadwick, F. M.; McKay, A. I. Adv. Organomet. Chem. 2016, 66, 223. 8. (a) Labinger, J. A.; Bercaw, J. E. Nature 2002, 417, 511; (b) Crabtree, R. H. J. Organomet. Chem. 2004, 689, 4083; (c) Wencel-Delord, J.; Dröge, T.; Liu, F.; Glorius, F. Chem. Soc. Rev. 2011, 40, 4740; (d) Jia, C.; Kitamura, T.; Fujiwara, Y. Acc. Chem. Res. 2001, 34, 633; (e) Crabtree, R. H. J. Chem. Soc. Dalton Trans. 2001, 2437. 9. (a) Caballero, A.; Despagnet-Ayoub, E.; Díaz-Requejo, M. M.; Díaz-Rodríguez, A.; González-Núñez, M. E.; Mello, R.; Muñoz, B. K.; Ojo, W.-S.; Asensio, G.; Etienne, M.; Pérez, P. J. Science 2011, 332, 835; (b) Cook, A. K.; Schimler, S. D.; Matzger, A. J.; Sanford, M. S. Science 2016, 351, 1421; (c) Smith, K. T.; Berritt, S.; González-Moreiras, M.; Ahn, S.; Smith, M. R., III; Baik, M.-H.; Mindiola, S. J. Science 2016, 351, 1424; (d) Periana, R. A.; Taube, D. J.; Gamble, S.; Taube, H.; Satoh, T.; Fujii, H. Science 1998, 280, 560; (e) Lin, M.; Sen, A. Nature 1994, 368, 613. 10. (a) Evans, D. R.; Drovetskaya, T.; Bau, R.; Reed, C. A.; Boyd, P. D. W. J. Am. Chem. Soc. 1997, 119, 3633; (b) Castro-Rodriguez, I.; Nakai, H.; Gantzel, P.; Zakharov, L. N.; Rheingold, A. L.; Meyer, K. J. Am. Chem. Soc. 2003, 125, 15734; (c) Andreychuk, N. R.; Emslie, D. J. H. Angew. Chem. Int. Ed. 2013, 52, 1696. 11. Ball, G. E. Spectrosc. Prop. Inorg. Organomet. Compd. 2010, 41, 262. 12. (a) Stolz, I. W.; Dobson, G. R.; Sheline, R. K. J. Am. Chem. Soc. 1962, 84, 3589; (b) Perutz, R. N.; Turner, J. J. J. Am. Chem. Soc. 1975, 97, 4791. 13. (a) Janowicz, A. H.; Bergman, R. G. J. Am. Chem. Soc. 1982, 104, 352; (b) Hoyano, J. K.; Graham, W. A. G. J. Am. Chem. Soc. 1982, 104, 3723. 14. Calladine, J. A.; Torres, O.; Anstey, M.; Ball, G. E.; Bergman, R. G.; Curley, J.; Duckett, S. B.; George, M. W.; Gilson, A. I.; Lawes, D. J.; Perutz, R. N.; Sun, X.-Z.; Vollhardt, K. P. C. Chem. Sci. 2010, 1, 622. 15. Kuimova, M. K.; Alsindi, W. Z.; Dyer, J.; Grills, D. C.; Jina, O. S.; Matousek, P.; Parker, A. W.; Portius, P.; Sun, X.-Z.; Towrie, M.; Wilson, C.; Yang, J.; George, M. W. Dalton Trans. 2003, 3996. 16. (a) Bernskoetter, W. H.; Schauer, C. K.; Goldberg, K. I.; Brookhart, M. Science 2009, 326, 553; (b) Walter, M. D.; White, P. S.; Schauer, C. K.; Brookhart, M. J. Am. Chem. Soc. 2013, 135, 15933. 17. (a) Pike, S. D.; Thompson, A. L.; Algarra, A. G.; Apperley, D. C.; Macgregor, S. A.; Weller, A. S. Science 2012, 337, 1648; (b) Pike, S. D.; Chadwick, F. M.; Rees, N. H.; Scott, M. P.; Weller, A. S.; Krämer, T.; Macgregor, S. A. J. Am. Chem. Soc. 2015, 137, 820; (c) Chadwick, F. M.; Rees, N. H.; Weller, A. S.; Krämer, T.; Iannuzzi, M.; Macgregor, S. A. Angew. Chem. Int. Ed. 2016, 55, 3677; (d) Martínez-Martínez, A. J.; Tegner, B. E.; McKay, A. I.; Bukvic, A. J.; Rees, N. H.; Tizzard, G. J.; Coles, S. J.; Warren, M. R.; Macgregor, S. A.; Weller, A. S. J. Am. Chem. Soc. 2018, 140, 14958; (e) Bukvic, A. J.; Georgiana Crivoi, D.; Garwood, H. G.; McKay, A. I.; Chen, T. T. D.; Martínez-Martínez, A. J.; Weller, A. S. Chem. Commun. 2020, 56, 4328; (f ) McKay, A. I.; Krämer, T.; Rees, N. H.; Thompson, A. L.; Christensen, K. E.; Macgregor, S. A.; Weller, A. S. Organometallics 2017, 36, 22; (g) Chadwick, F. M.; Krämer, T.; Gutmann, T.; Rees, N. H.; Thompson, A. L.; Edwards, A. J.; Buntkowsky, G.; Macgregor, S. A.; Weller, A. S. J. Am. Chem. Soc. 2016, 138, 13369; (h) Chadwick, F. M.; McKay, A. I.; Martinez-Martinez, A. J.; Rees, N. H.; Krämer, T.; Macgregor, S. A.; Weller, A. S. Chem. Sci. 2017, 8, 6014; (i) McKay, A. I.; Bukvic, A. J.; Tegner, B. E.; Burnage, A. L.; Martınez-Martınez, A. J.; Rees, N. H.; Macgregor, S. A.; Weller, A. S. J. Am. Chem. Soc. 2019, 141, 11700; (j) Martínez-Martínez, A. J.; Royle, C. G.; Furfari, S. K.; Suriye, K.; Weller, A. S. ACS Catal. 2020, 10, 1984. 18. Boyd, T. M.; Tegner, B. E.; Tizzard, G. J.; Martínez-Martínez, A. J.; Neale, S. E.; Hayward, M. A.; Coles, S. J.; Macgregor, S. A.; Weller, A. S. Angew. Chem. Int. Ed. 2020, 59, 6177. 19. (a) Lian, T.; Bromberg, S. E.; Yang, H.; Proulx, G.; Bergman, R. G.; Harris, C. B. J. Am. Chem. Soc. 1996, 118, 3769; (b) Bromberg, S. E.; Yang, H.; Asplund, M. C.; Lian, T.; McNamara, B. K.; Kotz, K. T.; Yeston, J. S.; Wilkens, M.; Frei, H.; Bergman, R. G.; Harris, C. B. Science 1997, 278, 260. 20. Bartlett, S. A.; Besley, N. A.; Dent, A. J.; Diaz-Moreno, S.; Evans, J.; Hamilton, M. L.; Hanson-Heine, M. W. D.; Horvath, R.; Manici, V.; Sun, X.-Z.; Towrie, M.; Wu, L.; Zhang, X.; George, M. W. J. Am. Chem. Soc. 2019, 141, 11471. 21. (a) Grills, D. C.; Childs, G. I.; George, M. W. Chem. Commun. 2000, 1841; (b) George, M. W.; Haward, M. T.; Hamley, P. A.; Hughes, C.; Johnson, F. P. A.; Popov, V. K.; Poliakoff, M. J. Am. Chem. Soc. 1993, 115, 2286; (c) Childs, G. I.; Grills, D. C.; Sun, X.-Z.; George, M. W. Pure Appl. Chem. 2001, 73, 443. 22. (a) Hodges, P. M.; Jackson, S. A.; Jacke, J.; Poliakoff, M.; Turner, J. J.; Grevels, F. W. J. Am. Chem. Soc. 1990, 112, 1234; (b) Stolz, I. W.; Dobson, G. R.; Sheline, R. K. J. Am. Chem. Soc. 1963, 85, 1013; (c) Nasielski, J.; Kirsch, P.; Wilputte-Steinert, L. J. Organomet. Chem. 1971, 29, 269; (d) Boylan, M. J.; Braterman, P. S.; Fullarton, A. J. Organomet. Chem. 1971, 31, C29; (e) Brown, C. E.; Ishikawa, Y.; Hackett, P. A.; Rayner, D. M. J. Am. Chem. Soc. 1990, 112, 2530. 23. Kazlauskas, R. J.; Wrighton, M. S. J. Am. Chem. Soc. 1982, 104, 6005. 24. (a) Rest, A. J.; Sodeau, J. R.; Taylor, D. J. J. Chem. Soc. Dalton Trans. 1978, 651; (b) Young, R. D.; Lawes, D. J.; Hill, A. F.; Ball, G. E. J. Am. Chem. Soc. 2012, 134, 8294. 25. Sawyer, K. R.; Cahoon, J. F.; Shanoski, J. E.; Glascoe, E. A.; Kling, M. F.; Schlegel, J. P.; Zoerb, M. C.; Hapke, M.; Hartwig, J. F.; Webster, C. E.; Harris, C. B. J. Am. Chem. Soc. 2010, 132, 1848.

520

Alkane s-Complexes

26. Brookhart, M.; Chandler, W.; Kessler, R. J.; Liu, Y.; Pienta, N. J.; Santini, C. C.; Hall, C.; Perutz, R. N.; Timney, J. A. J. Am. Chem. Soc. 1992, 114, 3802. 27. Belt, S. T.; Ryba, D. W.; Ford, P. C. Inorg. Chem. 1990, 29, 3633. 28. (a) Cowan, A. J.; Portius, P.; Kawanami, H. K.; Jina, O. S.; Grills, D. C.; Sun, X. Z.; McMaster, J.; George, M. W. Proc. Natl. Acad. Sci. 2007, 104, 6933; (b) Ball, G. E.; Brookes, C. M.; Cowan, A. J.; Darwish, T. A.; George, M. W.; Kawanami, H. K.; Portius, P.; Rourke, J. P. Proc. Natl. Acad. Sci. 2007, 104, 6927; (c) Childs, G. I.; Colley, C. S.; Dyer, J.; Grills, D. C.; Sun, X. –Z.; Yang, J.; George, M. W., J. Chem. Soc. Dalton Trans. 2000, 1901; (d) Creaven, B. S.; Dixon, A. J.; Kelly, J. M.; Long, C.; Poliakoff, M. Organometallics 1987, 6, 2600; (e) Sun, X.-Z.; Grills, D. C.; Nikiforov, S. M.; Poliakoff, M.; George, M. W. J. Am. Chem. Soc. 1997, 119, 7521. 29. (a) Calladine, J. A.; Duckett, S. B.; George, M. W.; Matthews, S. L.; Perutz, R. N.; Torres, O.; Vuong, K. Q. J. Am. Chem. Soc. 2011, 133, 2303; (b) Torres, O.; Calladine, J. A.; Duckett, S. B.; George, M. W.; Perutz, R. N. Chem. Sci. 2015, 6, 418. 30. Duckett, S. B.; George, M. W.; Jina, O. S.; Matthews, S. L.; Perutz, R. N.; Sun, X.–Z.; Vuong, K. Q., Chem. Commun. 2009, 1401. 31. Young, R. D.; Hill, A. F.; Hillier, W.; Ball, G. E. J. Am. Chem. Soc. 2011, 133, 13806. 32. Nayak, S. K.; Burkey, T. J. J. Am. Chem. Soc. 1993, 115, 6391. 33. (a) George, M. W.; Dougherty, T. P.; Heilweil, E. J. J. Phys. Chem. 1996, 100, 201; (b) Anfinrud, P. A.; Han, C.; Lian, T.; Hochstrasser, R. M. J. Phys. Chem. 1990, 94, 1180. 34. Poliakoff, M.; Turner, J. J. J. Chem. Soc. Dalton Trans. 1974, 2276. 35. Mawby, R. J.; Perutz, R. N.; Whittlesey, M. K. Organometallics 1995, 14, 3268. 36. (a) Blake, A. J.; George, M. W.; Hall, M. B.; McMaster, J.; Portius, P.; Sun, X.-Z.; Towrie, M.; Webster, C. E.; Wilson, C.; Zaric, S. D. Organometallics 2008, 27, 189; (b) Guan, J.; Wriglesworth, A.; Sun, X.-Z.; Brothers, E. N.; Zaric, S. D.; Evans, M. E.; Jones, W. D.; Towrie, M.; Hall, M. B.; George, M. W. J. Am. Chem. Soc. 2018, 140, 1842. 37. (a) Asbury, J. B.; Ghosh, H. N.; Yeston, J. S.; Bergman, R. G.; Lian, T. Organometallics 1998, 17, 3417; (b) George, M. W.; Hall, M. B.; Jina, O. S.; Portius, P.; Sun, X.-Z.; Towrie, M.; Wu, H.; Yang, X.; Zaric, S. D. Proc. Natl. Acad. Sci. 2010, 107, 20178; (c) George, M. W.; Hall, M. B.; Portius, P.; Renz, A. L.; Sun, X.-Z.; Towrie, M.; Yang, X. Dalton Trans. 2011, 40, 1751. 38. Bridgewater, J. S.; Netzel, T. L.; Schoonover, J. R.; Massick, S. M.; Ford, P. C. Inorg. Chem. 2001, 40, 1466. 39. Geftakis, S.; Ball, G. E. J. Am. Chem. Soc. 1998, 120, 9953. 40. Lawes, D. J.; Geftakis, S.; Ball, G. E. J. Am. Chem. Soc. 2005, 127, 4134. 41. Yau, H. M.; McKay, A. I.; Hesse, H.; Xu, R.; He, M.; Holt, C. E.; Ball, G. E. J. Am. Chem. Soc. 2016, 138, 281. 42. (a) Calvert, R. B.; Shapley, J. R. J. Am. Chem. Soc. 1978, 100, 7726; (b) Brookhart, M.; Green, M. L. H.; Wong, L. L. In Progress in Inorganic Chemistry; Lippard, S. J., Ed.; Wiley: New York, 1988; p 1; (c) Parkin, G. Acc. Chem. Res. 2009, 42, 315. 43. (a) Cordero, B.; Gómez, V.; Platero-Prats, A. E.; Revés, M.; Echeverría, J.; Cremades, E.; Barragán, F.; Alvarez, S. Dalton Trans. 2008, 2832–2838; (b) Pyykkö, P. J. Phys. Chem. A 2015, 119, 2326. 44. Alvarez, S. Dalton Trans. 2013, 42, 8617. 45. Bloch, E. D.; Queen, W. L.; Krishna, R.; Zadrozny, J. M.; Brown, C. M.; Long, J. R. Science 2012, 335, 1606. 46. (a) Wu, H.; Zhou, W.; Yildirim, T. J. Am. Chem. Soc. 2009, 131, 4995; (b) Hulvey, Z.; Vlaisavljevich, B.; Mason, J. A.; Tsivion, E.; Dougherty, T. P.; Bloch, E. D.; Head-Gordon, M.; Smit, B.; Long, J. R.; Brown, C. M. J. Am. Chem. Soc. 2015, 137, 10816. 47. (a) Lei, X.; Yang, J.; Lin, X.; Dai, Q.; Cheng, Q.; Guo, L.; Li, H. Chin. Sci. Bull. 2009, 54, 3244; (b) Pan, Q.; Guo, P.; Duan, J.; Cheng, Q.; Li, H. Chin. Sci. Bull. 2012, 57, 3867; (c) Hill, R. J.; Cranswick, L. M. D. J. Appl. Crystallogr. 1994, 27, 802. 48. (a) Balcells, D.; Clot, E.; Eisenstein, O. Chem. Rev. 2010, 110, 749; (b) Boutadla, Y.; Davies, D. L.; Macgregor, S. A.; Poblador-Bahamonde, A. I. Dalton Trans. 2009, 5820. 49. (a) Chan, B.; Ball, G. E. J. Chem. Theory Comput. 2013, 9, 2199; (b) Thenraj, M.; Samuelson, A. G. J. Comput. Chem. 2015, 36, 1818; (c) Cobar, E. A.; Khaliullin, R. Z.; Bergman, R. G.; Head-Gordon, M. Proc. Natl. Acad. Sci. USA 2007, 104, 6963; (d) Zaric, S.; Hall, M. B. J. Phys. Chem. A 1997, 101, 4646; (e) Cundari, T. R. Organometallics 1993, 12, 1998–2000. 50. Bergman, R. G.; Cundari, T. R.; Gillespie, A. M.; Gunnoe, T. B.; Harman, W. D.; Klinckman, T. R.; Temple, M. D.; White, D. P. Organometallics 2003, 22, 2331. 51. Lawes, D. J.; Darwish, T. A.; Clark, T.; Harper, J. B.; Ball, G. E. Angew. Chem. Int. Ed. 2006, 45, 4486. 52. (a) McNamara, B. K.; Yeston, J. S.; Bergman, R. G.; Moore, C. B. J. Am. Chem. Soc. 1999, 121, 6437; (b) Bengali, A. A.; Arndtsen, B. A.; Burger, P. M.; Schultz, R. H.; Weiller, B. H.; Kyle, K. R.; Moore, C. B.; Bergman, R. G. Pure Appl. Chem. 1995, 67, 281.

1.17

Dinitrogen Binding and Functionalization

Jeremy E Weber, Samuel M Bhutto, Alexandre T-Y Genoux, and Patrick L Holland, Department of Chemistry, Yale University, New Haven, CT, United States © 2022 Elsevier Ltd. All rights reserved.

1.17.1 Introduction and scope 1.17.2 Synthesis and characterization of N2 complexes 1.17.2.1 General procedures and experimental considerations 1.17.2.2 Key techniques and pitfalls to avoid 1.17.2.3 Synthesis of metal–N2 complexes 1.17.2.3.1 From atmospheric dinitrogen 1.17.2.3.2 From azides and hydrazines 1.17.2.3.3 From nitride coupling 1.17.3 N2 and its interactions with metals 1.17.3.1 Special properties of dinitrogen 1.17.3.2 Weakening of dinitrogen upon binding; comment on “activation” 1.17.3.3 Periodic trends and binding modes of N2 1.17.3.3.1 End-on terminal (Z1-N2) 1.17.3.3.2 End-on/end-on (m-Z1:Z1-N2) 1.17.3.3.3 Side-on (Z2-N2) 1.17.3.3.4 Side-on/side-on (m-Z2:Z2-N2) 1.17.3.3.5 End-on/side-on (m-Z1:Z2-N2) 1.17.3.3.6 N2 coordination to more than two metals 1.17.3.4 Supporting ligand environment 1.17.3.4.1 Steric effects 1.17.3.4.2 Ligand character 1.17.3.5 N2 splitting at transition metal complexes 1.17.3.5.1 N2 splitting to terminal nitrides 1.17.3.5.2 N2 splitting to bridging nitrides 1.17.4 Functionalization of N2 1.17.4.1 Forming NdH bonds at coordinated N2 1.17.4.1.1 Protonation of N2 1.17.4.1.2 NdH bonds from H2 1.17.4.1.3 Catalytic N2 reduction to NH3 1.17.4.1.4 NdH bond strengths 1.17.4.1.5 Concerted PCET 1.17.4.2 Forming other bonds at coordinated N2 1.17.4.2.1 NdSi bonds 1.17.4.2.2 NdC bonds 1.17.4.2.3 NdB bonds 1.17.4.2.4 Other N–X bonds 1.17.5 Summary and perspectives Acknowledgments References

1.17.1

521 522 522 523 523 523 523 524 525 525 525 526 527 527 528 530 531 531 532 532 533 534 534 535 536 536 536 537 539 540 541 542 542 545 549 550 550 550 550

Introduction and scope

Dinitrogen (N2) is ubiquitous, comprising more than 75% of the Earth’s atmosphere and more than 99% of the nitrogen atoms in the biosphere. This abundance is attributable to the thermodynamic stability and low reactivity of N2. For organometallic chemists, the most familiar application of N2 is also due to its low reactivity, as N2 is a common inert atmosphere for air-sensitive compounds. However, it is crucial that N2 can react, because the myriad applications of N-containing molecules are dependent on conversion of atmospheric N2 into more tractable forms (“nitrogen fixation”). Nitrogen fixation underlies major parts of the chemical industry, as it is responsible for the global production of millions of tonnes of ammonia (NH3) and nitric acid (HNO3).1,2 Studies of N2 reactivity have also gained popularity because nitrogen is an essential element for life, and nitrogen fixation is part of the web of reactions that control the biosphere.3 The fundamental challenge of nitrogen fixation is not thermodynamic: the standard potential for the reduction of N2 to NH3/ NH+4 is near that of the normal hydrogen electrode in both aqueous and organic solvents,4 and though the triple bond of N2 is very

Comprehensive Organometallic Chemistry IV

https://doi.org/10.1016/B978-0-12-820206-7.00068-8

521

522

Dinitrogen Binding and Functionalization

strong, its complete reduction yields six new bonds that can balance the energy lost during NdN cleavage. The real challenge with N2 reactions is kinetic, and high temperatures or extremely cathodic potentials are often used to overcome high activation barriers. Since binding of N2 to transition metals often increases its reactivity, organometallic chemists have striven to create metalN2 complexes as a first step toward catalysis. Chemists are also inspired by natural nitrogen fixation, which is performed by nitrogenase enzymes. The active sites of these enzymes are metal clusters that contain a carbide bridging between six iron atoms. Since the active sites for nitrogen fixation in nature contain FedC bonds, it is evident that nature is using organometallic chemistry for nitrogen fixation! Interestingly, the carbide is derived from S-adenosylmethionine,5 a sulfonium ion that has not been used by organometallic chemists, suggesting the potential for future research in bioorganometallic chemistry. Nitrogenases and synthetic complexes related to the enzymes are described in another chapter of Comprehensive Organometallic Chemistry IV, and therefore biological nitrogen fixation will not be discussed further in this chapter. Many of the compounds described in this chapter are not formally organometallic, as they lack metal-carbon bonds. However, the relationship of N2 binding to the binding of the isoelectronic and isolobal molecule CO, and the interest in N2 binding, weakening, cleavage, and functionalization within the community of organometallic chemists focusing on small-molecule activation, mean that the techniques, concepts, and interpretation of N2 reactions are often associated with organometallic chemistry. This chapter aims to provide a guide of both conventional wisdom and new developments in this area. Instead of attempting to exhaustively list all N2 complexes or reactions, we give representative examples, and we refer interested readers to numerous review articles and books.6–33

1.17.2

Synthesis and characterization of N2 complexes

1.17.2.1

General procedures and experimental considerations

Transition metal centers that bind N2 generally have low formal oxidation states, and are thus susceptible to oxidation and other reactions that preclude N2 binding. For example, ubiquitous O2 and H2O can oxidize low valent metals and form complexes that are inert to N2. Therefore, the study of N2 reactions typically includes careful measures to exclude O2 and H2O, starting with high purity reagents and a carefully purified atmosphere (glovebox or Schlenk technique). Even with these precautions, the complexes may lose N2 upon exposure to vacuum, and in many cases in situ spectroscopic methods are required for characterization. The choice of solvent can also introduce pitfalls. Easily reduced solvents like dichloromethane may be unstable to the reduced metal precursors. Solvents that are competent ligands (e.g. acetonitrile) may displace N2. It is more common to use solvents with low coordinating ability, such as alkanes, dialkyl ethers, and substituted arenes. The competition between solvent and N2 is metal-dependent, and thus metals capable of strong backbonding may bind N2 in tetrahydrofuran (THF) or in water, or ones with weak backbonding may bind N2 in benzene.34 Accurate measurement of N2 concentration in solution is imperative for the thermodynamic and kinetic analysis of any reaction of N2 with a metal complex. The solubility of N2 is solvent-dependent, and varies based on temperature and pressure. The mole fraction of N2 in some solvents at 1 atm has been tabulated over various temperatures,35,36 and these can be converted to concentration of N2 at saturation (Table 1). The solution N2 concentration can be controlled by adding a known volume of N2-saturated solvent, or by varying the pressure of N2 within a reaction vessel constructed from thick-walled glass or stainless steel. It is important to note that gas-liquid mixing is slow and the rate of dissolving a gas is proportional to the surface area of the gas-liquid interface. Therefore, vigorous bubbling or shaking is often required to increase the rate of mixing. Accelerating the reaction between N2 and unstable transient reduced species helps N2 binding to outcompete decomposition. The synthesis of N2 complexes can exploit the temperature dependence of the equilibrium constant for N2 binding, which increases with lower temperature. This temperature dependence is sometimes incorrectly attributed to the influence of the entropy term TDS, but according to the van’t Hoff equation, the temperature dependence of the equilibrium constant actually comes from the enthalpy term DH/RT. Therefore, it is more accurate to state that N2 binding is always entropically unfavorable, and thus N2 binding must be driven by a favorable enthalpy change which has a steep temperature dependence. Table 1

N2 concentration at saturation in some common solvents at 1 atm.

Solvent

T (K)

[N2] (mM)

n-Hexane

298 273 213 298 298 293 273 195 298

10.6 10.8 11.1 5.1 6.4 11.6 11.9 12.9 0.66

Toluene Tetrahydrofuran Diethyl ether

Water

Dinitrogen Binding and Functionalization

1.17.2.2

523

Key techniques and pitfalls to avoid

X-ray crystallography is the standard method for determining the connectivity of atoms in organometallic chemistry. However, the crystallographic characterization of N2 complexes presents several distinct challenges. The first is Fourier truncation effects, which in bridging N2 complexes can give artifacts near the N2 and inaccurate N atom positions.37 Secondly, libration (wagging of a multiatomic unit) often gives systematic errors in bond distances to the metal in terminal N2 complexes. Thirdly, the inherent random error in crystallography from measurement uncertainty and thermal motions is often comparable to the small variations in NdN distances. There are documented cases of errors in crystallographic NdN distances on the order of 0.1 A˚ , which serve as a warning.38 Finally, X-ray crystallography may not show a discernable peak for nearby H atoms. In order to definitively show that N2 is a component of a new compound, vibrational spectroscopy is recommended. Vibrational bands from NdN stretching modes have characteristic frequencies (typically 1700–2300 cm−1) that can be measured using infrared (IR) and resonance Raman (RR) spectroscopy, and verified using the shift in samples from 15N2. Sometimes the NdN stretch is not a localized normal mode, and in these cases the measured shift from 14N2 to 15N2 is less than that calculated from the harmonic oscillator model. In addition, there may be coupling between the NdN stretching band and overtones or combination bands (“Fermi resonance”), which leads to doubling or shifting of peaks.39 Often the product of interest during N2 reduction is NH3, but accurate detection of this compound has its own set of challenges. Quantitative measurement of NH3 can be performed using ion chromatography, 1H nuclear magnetic resonance (NMR) spectroscopy, or a colorimetric test.40–43 The most common test involves the oxidation of phenol with H2O2 in the presence of ammonia to form indophenol, whose blue color gives an absorption at 625 nm that can be quantitated using visible spectrophotometry.44 However, other oxidizable compounds can interfere with indophenol formation, so it is best to vacuum transfer volatile materials under basic conditions prior to analysis.45 Care must be taken when quantifying newly synthesized NH3 because it is present in the ambient atmosphere (even in human breath), and can adsorb to laboratory surfaces.46,47 Thus, it is imperative to quantitate background NH3 using negative controls. Labeled ammonia from 15N2 must use separate controls, because commercial sources of 15 N2 are often less pure than commercial N2.48,49 To remove impurities from 15N2, the gas can be passed through acidic and alkaline solutions before introduction into a reaction cell.

1.17.2.3 1.17.2.3.1

Synthesis of metal–N2 complexes From atmospheric dinitrogen

The most attractive source of N2 for use in synthetic chemistry is atmospheric N2. A typical approach for the synthesis of N2 complexes is to treat a halide precursor with a reductant such as an alkali metal (Fig. 1). Because these harsh reductants can give side reactions, it can help to perform the reaction at low temperature (this also increases the equilibrium constant for N2 binding as described above). In order to accelerate the reduction reaction relative to the decomposition of precursors, many practitioners have transitioned to using other alkali metal reductants. A classic example is sodium amalgam (Na/Hg), but concerns over the toxicity of mercury have largely caused this to be replaced by the use of alkali metals supported on graphite (KC8 is very common, and RbC8 and CsC8 are also known), or alkali metal salts of naphthalene or anthracene radical anions.38,50,51 The alkali metals on graphite have a high surface area which makes them react rapidly despite being insoluble. These compounds have a fixed stoichiometry with intercalated alkali metal cations between reduced graphene sheets. When reduction with an alkali metal is successful, the halide ligand is replaced by N2 (Fig. 1). Mechanistic studies indicate that reduction precedes dissociation of X−, and sometimes the corresponding anionic halide complex (or its alkali metal adduct) can be isolated under an argon atmosphere.52 Reductive elimination is an alternative approach for generating metal species in low oxidation states for N2 binding. An example is the cyclometallated iridium complex in Fig. 2, which reversibly binds N2 accompanied by formation of the CdH bond in the adamantyl group.53 Photochemical reductive elimination can also lead to N2 binding.54 Displacement of a neutral ligand is a third way to bring about N2 binding. Sometimes the equilibrium is driven toward the N2 complex because the metal center is strongly backbonding, for example the classic ruthenium complexes from Taube that are illustrated on the left of Fig. 3.55 In an intriguing recent example, displacement of the normally strong PMe3 ligand by N2 occurs during formation of the formally iron(IV) species on the right of Fig. 3; this was rationalized by the thin steric profile of N2 but may also be driven by greater backbonding in the neutral product.56

1.17.2.3.2

From azides and hydrazines

Hydrazine (N2H4) and diazene (N2H2) are unstable with respect to disproportionation to ammonia and N2. Thus, hydrazine is a feasible precursor to bound N2. This was the method used to prepare the first N2 complex of a transition metal in 1965.57

Mn+

X

N2, e– – X–

Fig. 1 General approach for N2 binding upon reduction of a metal halide complex.

M(n–1)+

N2

524

Dinitrogen Binding and Functionalization

Ir

P

iPr iPr

N2

H

– N2

H iPr iPr

iPr iPr

P

H

Ir

N

N

P iPr iPr

P

Fig. 2 Formation of an N2 complex by reversible reductive elimination.

+

H3N H3N

NH3 Ru NH3

2+ NH3

N2

H3N

NH3

NH3

H3N

NH3 Ru N2

R

2+

N

NH3

Fe

NH3

Me3P

PMe3 iPr

PMe3

N

R–, N2

Fe

– PMe3

Me3P N

R = Me, Bn

PMe3 iPr

N

Fig. 3 Examples of displacement of a neutral ligand by N2.

Cl

PPh3 Cl Re O Ph3P Cl

O +

N H

Ph

NH2

PPh3 HCl, EtOH

Ph3P N N PMe2Ph Cl Re C Cl O Ph benzene/MeOH Ph3P

PhMe2P PhMe2P

N N Re Cl

PMe2Ph PMe2Ph

Fig. 4 Synthesis of an N2 complex from an acyl hydrazine.

A less common route uses acyl azides and hydrazines. Treating Vaska’s complex, trans-IrCl(CO)(PPh3)2, with an acyl azide (ArC(O)N3) leads to the replacement of the CO ligand with N2, forming trans-IrCl(N2)(PPh3)2.58–61 This reaction may involve an acyl isocyanate (ArC(O)NCO) intermediate that reacts with exogenous alcohol or water to generate the open site for N2. Aryl- and acylhydrazines can also be used as precursors to N2 complexes.62–64 For example, treating ReOCl3(PPh3)2 with HCl, PPh3, and 1,2-benzoylhydrazine leads to the rhenium(III) benzoyldiazenido complex, ReCl2(PPh3)2(N2C(O)Ph). When this compound is treated with MeOH and phosphine, the benzoyl group is lost as methylbenzoate and it is possible to isolate the terminal N2 complex trans-Re(N2)(PMe2Ph)4Cl (Fig. 4). In both cases, the azide and hydrazine are proposed to be the source of the N2 ligand.

1.17.2.3.3

From nitride coupling

Nitride coupling is another route to M–N2 bonds, and has been observed in a number of late-metal systems.65–69 This strategy has recently gained increasing attention in the context of catalytic NH3 oxidation to N2.70 This route is the microscopic reverse of N2 splitting (see Section 1.17.3.5), and involves the formal reduction of each metal by three electrons. Therefore, it is observed only for high-valent nitride complexes, and requires a relatively weak M–N multiple bond. A recent instructive example uses an Ir pincer system (PNP ¼ N(CHCHPtBu2)2, Fig. 5).71,72 Photolysis of the parent azide, (PNP)IrN3, or chemical oxidation of the azide to form [(PNP)IrN]+ followed by reduction, leads to a (PNP)IrN radical (the unpaired electron lies in an iridium-nitrogen p orbital, as indicated by a dot above the bond in Fig. 5). This metalloradical can undergo homocoupling to form a new NdN bond.

PtBu2 N Ir N3 PtBu2

[FeCp2][PF6] DCM –N2

hQ, THF –N2

PtBu2 N Ir N PtBu2

CoCp*2 THF PtBu2 PF6 N Ir N PtBu2

Fig. 5 Example of forming an N2 complex through nitride coupling.

THF

t PtBu2 Bu2P

0.5

N Ir PtBu2

N N tBu

Ir N 2P

Dinitrogen Binding and Functionalization

1.17.3

N2 and its interactions with metals

1.17.3.1

Special properties of dinitrogen

525

In organometallic chemistry, chemists often draw parallels between N2 and the more familiar organometallic diatomic ligand, carbon monoxide. This analogy is compelling since the N2 and CO molecules are isoelectronic and isolobal, and both interact with metals through s-donation of a lone pair and through backbonding into unoccupied p orbitals (Fig. 6).73 However, it is important to notice some important differences between these molecules and their interactions with metals. N2 is appreciably less basic, with a proton affinity (494 kJ mol−1) that is much smaller than that for the carbon atom of CO (594 kJ mol−1).17,74 In addition, the p orbitals of CO are polarized toward the carbon atom, making it a much stronger p-acceptor than N2 where the p orbitals are unpolarized. For these reasons, N2 binds to metals 100 kJ mol−1 less strongly than CO in common zerovalent complexes.75

Fig. 6 Molecular orbital interactions in end-on N2 complexes.

Other properties of N2 are exceptional as well. Namely, N2 has a large ionization energy (15.6 eV), poor electron affinity (−1.8 eV) and small proton affinity (494 kJ mol−1), making it resistant to charge transfer or protonation pathways.17,74 This relative stability contrasts with the reactivity of other diatomics (like NO and O2) where charge transfer facilitates binding and reaction. Further, N2 lies in a deep trough in the Frost diagram of nitrogen compounds, meaning that it is difficult to reduce or oxidize by 1–2 electrons.76 Part of the difficulty is that the two-electron reduction product of N2, diazene (N2H2), has a double bond that is weakened by repulsion between the N lone pairs. As a consequence, the formal bond dissociation energy of the first p bond in N2 is much greater than the other two bonds. Thus, DH for the reaction of N2 and H2 to form N2H2 is more than 200 kJ mol−1 uphill, whereas the reaction of CO and H2 to form H2CO is roughly thermoneutral.17,74 Because of these thermodynamic considerations, it is difficult for the energetic cost for N2 reduction to be repaid until later steps in the reductive pathway. The high energy of partially-reduced intermediates with N2 distinguishes it from alkynes and alkenes (where each hydrogenation step is exothermic).

1.17.3.2

Weakening of dinitrogen upon binding; comment on “activation”

Upon binding to a transition metal, chemists typically try to gauge N2 “activation.” The term “activation” is unfortunate, because it is used for both ground-state properties and reactivity, and often leads chemists to conflate the two. However, the ground-state properties of the bound N2 (NdN bond length, force constant, stretching frequency) are often uncorrelated to the reactivity of the bound N2. In this chapter, we will distinguish between ground-state and reactivity effects through the terms weakening and activation, respectively. Activation is delineated in this way to draw analogies with other chemistries (such as CdH activation meaning the cleavage of the bond). Additionally, in the literature the term “activation” may be used to mean “weakening.” This creates a precarious situation because even N2 ligands that seem to be weakened/reduced substantially through binding can remain unreactive toward other reagents, and N2 only dissociates to return the electron density to the metal(s). Conversely, functionalization of N2 is often observed in complexes with little ground-state weakening of the NdN bond, indicating that weakening is not necessary for activation. Here we will discuss measures of weakening, while the reactions of N2 will be discussed below. Weakening of the NdN bond may be quantified in several ways, each of which has advantages and disadvantages. Therefore, multiple forms of characterization should be used when assessing how much N2 has been weakened by binding to the metal(s). The most commonly used measure is the NdN distance, which is typically ascertained using an X-ray crystallographic structure. It is inadvisable to attach great significance to differences in NdN bond lengths that are less than 0.01–0.02 A˚ (see Section 1.17.2.2). Another way to quantify NdN bond weakening is through the NdN stretching frequency. This method is less prone to random errors, as the range of NdN stretching frequencies varies over hundreds of wavenumbers while measurements have precision better than 2 cm−1. The stretching frequencies correlate with the NdN bond lengths through the empirical “Badger’s Rule.”9,77 Terminal N2 ligands give intense bands in IR spectra, but when N2 lies on a center of symmetry, the NdN stretching band is not IR-active, and it is necessary to utilize Raman spectroscopy, which sometimes does not yield useful spectra because the laser causes dissociation of N2 or degradation of the complex.

526

Dinitrogen Binding and Functionalization

Table 2 Benchmark NdN bond distances and stretching frequencies for N2 weakening.78

Free N2 [N^N]0 Free N2H2 [N ¼N]2− [N ¼N]3− Free N2H4 [NdN]4−

˚) N–N Distance (A

Stretching Frequency (cm−1)

1.10 1.1–1.2 1.25 1.25–1.35 1.40 1.45 1.45–1.60

2331 1700–2331 1583/1529 1200–1700 989–1040 885 700–1100

The extent of N2 weakening is influenced by the characteristics of the metal site, for example its oxidation state, electronegativity, atomic radius, and geometry. Benchmarks for comparing bond length and stretching frequency are free N2 (triple bond), diazene (N2H2, double bond) and hydrazine (N2H4, single bond). In complexes of N2, reduction to the diazene level is described as N2− 2 and to the hydrazine level as N4− 2 (Table 2). This formal electron transfer from the metal to the ligand is a form of redox-noninnocence. Since the strength of the interaction between a transition metal and N2 is largely dependent on backbonding or electron transfer from metal d orbitals, N2 binding typically requires a metal precursor with greater than zero d electrons. These precursors are known across the transition series as well as the lanthanides and actinides. With early transition metal ions (groups 3–5; lanthanides; actinides), N2 binding is most often accomplished by a formal reduction of the N2 ligand.79–84 Considered in this way, formal electron transfer to N2 gives a d0 metal center which electrostatically stabilizes the Nn– 2 (n ¼ 1-3) product. With later transition metals (groups 6–9), N2 complexes are most abundant, while there are few examples of metalN2 complexes in group 10 and later.22,85 Group 11 N2 complexes are rare, which can be attributed to the lesser backbonding ability of these electronegative metals, and no group 12 N2 complexes are yet known.86,87 Recently, it has been shown that boron and calcium complexes of N2 can be isolated, and this occurs through full electron transfer from formally boron(II) or calcium(I) intermediates to reduce N2.88,89

1.17.3.3

Periodic trends and binding modes of N2

Dinitrogen can interact with one or more transition metals, through a number of binding modes (Fig. 7). Structural analyses and correlations with reactivity have been reviewed.20,90 Important aspects are summarized in different sections for each binding mode. Diagrams of the periodic table use the darkness of the color to indicate increasing numbers of crystallographically characterized complexes with that binding mode, using the Cambridge Crystallographic Structure Database.91

Fig. 7 Binding modes of N2. In this figure, the NdN bonds are generally drawn as triple bonds, but the bond order may be reduced in the complexes as described below.

Dinitrogen Binding and Functionalization

1.17.3.3.1

527

End-on terminal (
1-N2 , Fig. 8)

The most common transition metal-N2 complexes have an end-on terminal binding mode which is denoted as Z1-N2. Though the analogy to CO would lead one to predict longer NdN bonds with more electropositive metals at the left of the transition series (higher energy d orbitals), terminal M–N2 complexes have similar NdN distances of 1.11  0.02 A˚ through the range from group 4 to group 10.91 This indicates that backbonding into terminal N2 is relatively weak, and the NdN bond is best described as a triple bond when bound in an end-on fashion.

Fig. 8 Z1-N2 complexes.

Metal complexes with N2 in a terminal binding mode are key intermediates for many of the well-defined catalysts that can activate and functionalize N2 (see Section 1.17.4.1.3). Binding to a transition metal polarizes the N2, and the distal N atom becomes nucleophilic and can bind to Lewis acids.92,93 Lewis acid binding weakens the NdN bond by helping to withdraw electron density from the metal into the N2 unit (Fig. 9). Often these Lewis acids contain main-group metals or metalloids, and the examples of alkali metal interactions with N2 are particularly numerous.94 Systematic analyses of different alkali metals, boranes, and H-bond donors were recently reported with Re and Fe complexes, showing that N2 could be cooperatively weakened by the transition metal and the Lewis acids as indicated by decreases of the NdN stretching frequency by up to 172 cm−1.95,96 Lewis acid interactions also increase the ability to protonate the distal nitrogen, even though one might have expected the Lewis acid to satiate the partial negative charge on this N atom. This ability of the Lewis acid to withdraw charge into the N2 unit can aid in subsequent N2 functionalization (see Fig. 33 below).

LA M

N

N

Fig. 9 A Lewis acid on the distal N of an end-on complex creates an additional dipole that enhances backbonding and N2 reactivity.

1.17.3.3.2

End-on/end-on (m-
1:
1-N2 , Fig. 10)

The nucleophilic distal N atom of an Z1-N2 complex can bind to a second transition metal. When this additional metal center has accessible d orbitals for backbonding, a linear MNNM structure is formed (Fig. 11). The cooperative backbonding from the two metals is much stronger than in mononuclear Z1-N2 complexes, because it allows electronic delocalization. For example, the formally vanadium(III) complex at the left of Fig. 11 has a longer and weaker NdN bond, and based on the distance is described as 4+ 97,99 V4+(N2− 2 )V . A series of iron, cobalt, and nickel b-diketiminate complexes with MNNM cores illustrate some of the periodic trends associated with this bridging mode.100,101 Moving to the more electronegative elements to the right of the periodic table, there is less weakening of the NdN bond (Fe: 1.19 A˚ /1778 cm−1; Co: 1.14 A˚ ; Ni: 1.12 A˚ /2164 cm−1). Further to the right, the analogous copper(I) complex does not bind N2, although it can be induced to do so by including a third copper(I) ion in a constrained framework.86 In a separate example, the diiridium(I) complex at the right of Fig. 11 has almost no NdN weakening.98

528

Dinitrogen Binding and Functionalization

Fig. 10 m-Z1:Z1-N2 complexes.

Ar Ar

Ar Ar

V

N

N

V

Ar

Ar

Ar = mesityl d(N–N) = 1.233(8) Å

O

2–

PtBu2 Ir

Ar O

tBu

2P

N N

PtBu2

O Ar

Ir tBu

2P

O

Ar = 3,5-bis-(trifluoromethyl)phenyl d(N–N) = 1.119(6) Å

Fig. 11 Representative examples of end-on/end-on N2 complexes.97,98 Note that different resonance structures are used to suggest the NdN distances, though the actual bond order is often unknown.

The selectivity for terminal vs. bridging N2 is influenced most obviously by steric effects, and the use of very bulky ligands can prevent the formation of the bimetallic core.20 However, bridging does not occur in some systems despite small supporting ligands. This has been explained by Hasanayn and coworkers in terms of the number of electrons in MNNM p-orbitals; for example, octahedral dn complexes with n ¼ 1–5 favor bridging N2, but in octahedral d6 systems an antibonding MNNM molecular orbital is occupied and therefore the bridging mode is higher in energy.102 Despite the greater formal NdN bond reduction compared to terminally bound N2 complexes, the N2 unit in m-Z1:Z1-N2 complexes is not nucleophilic and rarely reacts with other reagents. This lack of reactivity may be due to the need to rehybridize at N2 for reaction with an electrophile, which would require significant electronic redistribution and geometric distortion. The lack of a nucleophilic site in the symmetric bridging N2 species distinguishes these from the Lewis acid adducts in Section 1.17.3.3.1 above, in which the N2 can be bent, and cooperation of the transition metal and the main-group metal heightens the reactivity of N2. Recent efforts have yielded a new type of non-linear FeNNFe bridge, which was motivated by surface studies indicating that the immediate precursor to N2 cleavage has the NdN bond roughly parallel to the surface.103–106 In these complexes, the FeNNFe cores are distorted away from linearity by ligands that are linked by organic groups or cation-p interactions.107–109 It is not yet clear whether this bending leads to greater reactivity as would be expected from the rehybridization at N. The common bridging mode in CO complexes is m-1,1, which has never been observed in N2 complexes, and this is another significant difference between N2 and CO.

1.17.3.3.3

Side-on (
2-N2)

To date, there are no isolated examples of mononuclear complexes with an N2 ligand in the side-on binding mode. However, this is a compelling synthetic target because computations suggest that side-on binding weakens the NdN bond more than end-on binding.101 The orbital interactions underlying this coordination mode are distinct from end-on binding. Here, the s interaction is comprised of the empty metal d orbital and filled p orbital of N2 while the p backbonding has different overlap with the empty p orbital of N2 (Fig. 12). Even though none have been isolated, a few studies have provided evidence for the transient existence of side-on N2 ligands (Fig. 13). In one case, singly-labeled complexes CpRe(CO)(L)(15N14N) (Cp ¼ cyclopentadienyl; L ¼ CO, PMe3, or P(OMe)3), which contain an 15N label specifically on the proximal position, convert to the isotopologue with the 15N in the distal position

Dinitrogen Binding and Functionalization

529

Fig. 12 Potential molecular orbital interactions for side-on N2 complexes.

Fig. 13 Evidence for transient side-on N2 species, which rapidly return to the more stable end-on binding mode.

without release of N2.110 This rotation of coordinated N2 implicates an Z2-N2 intermediate, though the lifetime may be very short (or it may be only a transition state with no lifetime). A more direct observation of side-on N2 was obtained from spectroscopic and crystallographic studies on single crystals of [Os(NH3)5(Z1-N2)]2+ during cryoscopic UV irradiation.111 Irradiation led to a new IR band at 1838 cm−1 accompanied by a component in the X-ray crystal structure indicating 5% population of [Os(NH3)5(Z2-N2)] [PF6]2. Later work showed that an analogous species can be generated with ruthenium analogues, and also used differential scanning calorimetry to show that the metastable side-on N2 species releases 4 kcal mol−1 of heat upon relaxation to the ground state.112 These studies indicate that side-on N2 binding is possible, but it has not yet been possible to test the reactivity of N2 in this ephemeral binding mode.

Fig. 14 . m-Z2:Z2-N2 Complexes.

530

Dinitrogen Binding and Functionalization

1.17.3.3.4

Side-on/side-on (m-
2:
2-N2 , Fig. 14)

Another bimetallic N2 binding mode is side-on/side-on (m-Z2:Z2-N2), which is commonly observed with early transition metals.79–84 There are two classes of side-on/side-on N2 compounds, one with more ionic interactions between the metal and the N2 unit, and one with more covalent interactions. Examples of the former class are shown at the left and center of Fig. 15.113–115 These are also observed with lanthanides, with good examples being the series of complexes [((Me3Si)2N)2(THF)2M]2(m-Z2:Z2-N2) 116,117 (M ¼ lanthanide) which have NdN bond lengths and stretching frequencies indicating N2− It is possible to reduce the 2 . 2 2 related compound [((Me3Si)2N)2(THF)Y]2(m-Z :Z -N2) by one electron to give [K(THF)6][((Me3Si)2N)2(THF)Y]2(m-Z2:Z2-N2) 118 (Fig. 16). Its longer NdN bond length of 1.401(6) A˚ led to the assignment of the bridging ligand as N3− 2 .

Ph nBu

Cp*

NiPr

iPrN

N

Cp*

Y Y N Cp* Cp* d(N–N) = 1.172(6) Å

Sc

Cp*

N

N

Cp* Sc

N

Me2Si

N

NiPr

iPrN

Ph

Ph

N Zr

Me2Si

nBu

d(N–N) = 1.250(1) Å

Zr

N

P Ph

SiMe2

P N

SiMe2

N

Ph

d(N–N) = 1.503(3) Å

Ph

Fig. 15 Examples of side-on/side-on bridging N2.

K(THF)6 THF N L Y 2YL3 + 2K N –2 KL L N2

2

L Y

L

KC8

THF

THF

L

N L Y N THF

L

L

L Y N N THF

THF

Y

K

L

L Y

THF L

2

Fig. 16 Various oxidation levels of m-Z :Z -N2 ligands. L ¼ N(SiMe3)2.

The other class of side-on/side-on bridging complexes, for example the compound at the right of Fig. 15, can be viewed as having stronger covalent interactions between the p-orbitals of N2 and the metals, in analogy to the Dewar-Chatt-Duncanson model. They may be compared to side-on/side-on O2 complexes that are well-known in copper chemistry.119 These side-on/side-on 120,121 N2 complexes are typically best described as having an N4− 2 bridge, which is nucleophilic. The preference for side-on/side-on vs. end-on/end-on coordination may be dictated by the energetics of the bonding d orbitals and by the steric accessibility of open coordination sites. In some instances, the N2 ligand can switch between side-on/side-on and end-on/ end-on. The influence of steric bulk on hapticity of N2 is illustrated by the group 4 complexes Cp 2M–N2–MCp 2 (Cp ¼ pentamethylcyclopentadienyl) and their close analogues. For Ti complexes, removal of three Me groups from one Cp 2M fragment and two Me groups from the other induces a hapticity change from end-on to side-on.122–124 For the larger Zr and Hf, removal of only one Me group from each Cp induces a change from end-on/end-on to side-on/side-on (Fig. 17).81,125,126 This trend arises because the sideon/side-on coordination mode requires close M–M distances which are only achievable with smaller ligands.

M N N M

M

N N

M

Fig. 17 Removal of one methyl group from each Cp ligand leads to a change from end-on to side-on binding. M ¼ Zr, Hf.

The oxidation state also influences the binding mode as shown by the N2-bridged Zr complex supported by an adamantylsubstituted ansa-bis(cyclopentadienyl) ligand.127 Oxidation and reduction of this complex alters the binding mode of the N2 ligand (Fig. 18). A change in stretching frequency from 1445 cm−1 to 1727 cm−1 indicates that the NdN bond is significantly strengthened by the oxidation and change from end-on/end-on to side-on/side-on.

Et2O

OEt2

Ad Cl Cl Zr N N Zr K

Me2Si

Ad

AgBPh4 SiMe2

Fig. 18 Change of N2 binding mode with oxidation state.

KC8

Ad

Cl Zr Me2Si

N

Ad

N

SiMe2 Zr Cl

+ KBPh4

Dinitrogen Binding and Functionalization

531

End-on/side-on (m-
1:
2-N2)

1.17.3.3.5

The rarest of the bimetallic coordination modes of N2 is end-on/side-on (m-Z1:Z2-N2, Fig. 19), which is known only for a few Ti, Zr, and Ta complexes.128–133 In each case, a short M–M distance is enforced by either a bridging ligand (H− or Cl−) or a fused Cp ligand, suggesting that the ligand constraint may be necessary to prevent N2 from adopting a more stable bridging mode. In this coordination mode, NdN distances are 1.19 A˚ to 1.31 A˚ indicating significant N2 weakening. The asymmetric interactions of the metals with N2 polarize the N2 ligand, which leads to facile reaction with electrophiles (see Fig. 53, for example).

R2

Ph

Ph Ph

N

Me2Si

N

H H Zr

Zr P N Ph

Me2Si

N

N

Ph

N

SiMe2

P Ph

R1

Cl Zr

Zr

SiMe2

N

N

R1 = R2 = SiMe3 R1 = iPr, R2 = Me R1 = tBu, R2 = Me Fig. 19 Two examples of end-on/side-on N2 complexes of Zr.

Ph Ti Ti N N Ti

N

Ti

Ti

N

Ph

N N N Cu N Cu N Cu N N N

Ti

Ph

Ti

N N N Sm N N Sm Sm N NN NN Sm

Ph tBu

P

P

Au P Au Au P

O

Au N

N

Au P Au P

P = PiPrPh2

O O THF Sm O N O O Sm O N Sm O O O O O

Na

O O

O = O tBu

Ph Ph

Ph

Ph (THF omitted)

OMe OMe MeO OMe

tBu

tBu

Fig. 20 Examples of complexes with N2 bridging between more than two metals.

1.17.3.3.6

N2 coordination to more than two metals

Trimetallic and higher nuclearity N2 complexes have been isolated in a few cases (Fig. 20). The side-on/end-on/end-on coordination mode has been identified in complexes of Ti, Co, and Cu.86,124,134,135 The Ti complexes are related to their bimetallic analogues and exhibit greatly weakened NdN bonds with an NdN distance of around 1.3 A˚ and an NdN stretching frequency of 1282 cm−1. In contrast, the Cu3 complex exhibits little NdN weakening (dNN ¼ 1.08–1.10 A˚ and nNN ¼ 1952 cm−1), and is held by a cage consisting of three tethered b-diketiminate units. The Co3 analogue of this compound was very recently reported, and a combination of computational and spectroscopic studies indicated more weakening (NdN stretching vibrations at 1717 and 1752 cm−1).135

532

Dinitrogen Binding and Functionalization

Other binding modes give NdN single bonds. The only metal to be isolated with side-on/side-on/side-on and end-on/side-on/ side-on/end-on coordination is Sm, and gives long NdN distances (1.4 A˚ or more).136–139 Finally, there exists one example of an N2 ligand bound to six Au centers with an NdN single bond distance of 1.42 A˚ .140 This complex is derived from hydrazine rather than N2, and has a closer analogy to doubly protonated hydrazine (N2H2+ 6 ).

1.17.3.4

Supporting ligand environment

Many supporting ligands have been used for N2 complexes. This section describes representative examples in which systematic changes in the supporting ligand have been studied, to give insight into the ways in which the organometallic chemist may modulate the binding of N2.

1.17.3.4.1

Steric effects

The isolation of N2 complexes typically requires the use of bulky ligands that protect the easily-disrupted bond between the metal and N2. Without steric bulk, many tridentate, tetradentate, and pentadentate supporting ligands could form 2:1 complexes, or could utilize bridging groups (e.g. halides, alkoxides, thiolates, carboxylates) to fill the open coordination site rather than N2 binding. Bulky supporting ligands destabilize these undesired products, and the slender N2 can coordinate in the small pockets that remain. A seminal example of bulky ligand design is triamidomolybdenum(III) complexes. Molybdenum complexes with smaller amido ligands, when reduced to molybdenum(III), had formed ModMo triple bonds or had bridging amido groups.141 However, by using amido groups with large t-butyl and xylyl substituents (Fig. 21), Cummins and coworkers were able to avoid these pathways and coordinate N2 between two Mo ions at low temperature (followed by NdN cleavage, as described below).142–144

Fig. 21 Sterically bulky ligands prevent ModMo bond formation and enable N2 binding.

Chirik and coworkers have studied the systematic variation of substituents on Cp rings in a series of group 4 metallocenes, which demonstrates the changes in N2 binding that can result from subtle steric modification. With Ti, they prepared bis(cyclopentadienyl)titanium(II) complexes where the cyclopentadienyl groups bear four methyl groups and a fifth group R ¼ Me, Et, iPr, tBu, Ph, 3,5-Xyl, SiMe3, SiMe2Ph, or SiMet2Bu.145 The larger substituents led to weaker N2 binding in the adducts Cp’2Ti(N2)2 (Cp’ ¼ Me4CpR), and the NdN stretching frequencies were high, both of which indicate that s-effects predominate over p-effects in this system (little backbonding). In the Ti system, the smallest substituents also led to formation of bridging N2 complexes, which were not observed with larger substituents. The analogous Zr systems showed a more marked change of coordination with a change in size of the substituents on Cp rings.146 With two Cp ligands on each formally zirconium(II) center, N2 binding gave three N2 per two Zr atoms including a m-Z1:Z1 bridging of one N2. With removal of a single methyl group from each Cp , the tetramethylcyclopentadienyl-supported complex gave side-on bridged N2 with much more extensive NdN bond weakening. These trends are understandable since the metal-metal distance is much shorter for m-Z2:Z2 bridging than m-Z1:Z1 bridging, and bulky ligands destabilize the former. Another demonstration of the influence of bulky ligands on N2 binding and weakening comes with iron b-diketiminate complexes (Fig. 22). In these systems, the N2 binding was assessed through reduction of an iron(II) chloride complex, which reduces the metal to iron(I) and creates an open coordination site on the iron for N2 binding (see Fig. 1 above).147 The bulkiest supporting ligands, with isopropyl groups in the ortho-positions of the aromatic rings, gave m-Z1:Z1 bridging N2.52,148 Reducing the size of the ortho-substituents to ethyl resulted in a system in which N2 binding did not occur, but instead the aryl rings bridged to a second metal.149 Further reduction in the size of the ligand gave a bis(nitride)tetrairon product that has an Fe3N2 core derived from N2 cleavage, implicating an intermediate with end-on/side-on/side-on N2 in a transient intermediate, as surmised from density functional theory (DFT) calculations on a truncated model.150,151 These examples show the great sensitivity of N2 binding mode and reactivity to the steric details of the ligands.

Dinitrogen Binding and Functionalization

533

Fig. 22 Steric variation of b-diketiminate complexes of iron causes substantial differences in N2 binding or cleaving ability.

1.17.3.4.2

Ligand character

Backbonding is important in most N2 complexes, so in general N2 binds more strongly and is more weakened when the ligand environment increases the electron density at the metal center. Thus, stronger s- and p-donors are often used as supporting ligands, and amido and phosphino donors are particularly common. A growing number of supporting ligands incorporate N-heterocyclic carbenes, in keeping with their strong s-donor capacity and their popularity in new supporting ligand designs. With early metals and lanthanides, there are numerous examples with alkoxo and amido supporting ligands. With the middle transition metals (groups 6–8), the requirements for bulkiness and for p-donor capacity are relaxed, and even small ligands and p-acceptor ligands like CO are found in some N2 complexes.152–157 In these d6 complexes, there is usually little weakening of the NdN bond. One example of a study with systematic variation of ligand p-donor properties uses (PNP)Ru(H)N2 complexes, where PNP represents a tridentate pincer ligand with two di(t-butyl)phosphino arms and a central N-donor (Fig. 23).158 The more strongly p-donating amido donors gave lower NdN stretching frequencies (though no changes in the NdN bond distances were statistically significant, as noted in Section 1.17.2.2 above). In addition, when the ligand was oxidized or protonated, the reduced p-donor character gave higher NdN stretching frequencies, characteristic of weaker backbonding.

strongest backbonding

weaker S donor: weaker backbonding

PtBu2

PtBu2

PtBu2

N Ru N N H PtBu 2

N Ru N N H t P Bu2

N Ru N N H PtBu

cationic: weaker backbonding PtBu2 H N Ru N N H PtBu 2

+

2

S acceptor: weaker backbonding PtBu2 N Ru N N H PtBu 2

Fig. 23 Electronic effects on the extent of backbonding in a series of ruthenium(II)dN2 complexes.

+

PtBu2 N Ru N N H PtBu 2

+

534

Dinitrogen Binding and Functionalization

The higher electronegativity of late transition metals is commonly offset by the use of powerful s-donor/p-donor ligands in order to encourage N2 binding. For example, Peters and coworkers have utilized tripodal phosphine ligands with different apical coordinating atoms that vary from X to L to Z type donors (silyl, alkyl, amine, phosphine, borane; Fig. 24), which bind N2 trans to the variable donor.45,159–162 In this series, the NdN stretching frequencies vary from 1905 to 2011 cm−1, with CP3 and BP3 giving the greatest and least N2 weakening, respectively. Interestingly, the complex with the boron donor, which has the strongest NdN bond, is the most catalytically active for reducing N2 to NH3.161,162 The discrepancy between ground-state properties and reactivity is a theme that carries through this chapter.

iPr iPr

2P

2P

n

N N Fe

PiPr

2

E

E = B, C, Al, Ga, Si

E

n

d(N–N) (Å)

B Al Ga Si C

0 0 0 –1 –1

1.06 1.12 1.13 -

νNN (cm–1) 2011 2003 1997 1920 1905

Fig. 24 Trends in ground-state N2 weakening with variation of the trans ligand in EP3 complexes.

1.17.3.5 1.17.3.5.1

N2 splitting at transition metal complexes N2 splitting to terminal nitrides

Certain low-valent complexes with d3 to d5 electron configurations are capable of splitting N2 homolytically to give two equivalents of a terminal nitride complex.163 Examples are shown in Fig. 25.78 The resultant terminal nitride complexes are formally oxidized by three electrons at each metal, which provides the six electrons required to form the two nitrides. This is the reverse of the nitride coupling pathway to N2 complexes mentioned above in Section 1.17.2.3.3.

Fig. 25 Examples of complexes that can split N2 through a bimetallic pathway to give terminal nitrides. Ar indicates the aryl groups shown above in Fig. 21.

N2 splitting reactions of this type are thought to begin with binding of N2 in an end-on/end-on bridging mode. The feasibility of subsequent NdN splitting is understood through the electronic structure of the intermediate MNNM species. One thermodynamic criterion for splitting is the presence of 10 electrons in the p system of this MNNM core, which provides the electrons for M–N triple bonds in the products.164 In addition, each of the successful systems lacks a ligand trans to the N2 binding site, which enhances the strength of the metal-nitride bond in the product.165 In this way, the N2 cleaving reaction can be exergonic despite the formidable NdN bond energy of N2.

Dinitrogen Binding and Functionalization

535

Despite the linear geometry of the MNNM unit, the transition state for NdN cleavage has a “zigzag” structure, because of symmetry constraints on orbital overlap that have been discussed in some detail in the literature (Fig. 26). In a few cases, the N2 splitting occurs photochemically but not thermally, but there is not yet understanding of the excited states that are effective for enabling the photochemical reaction.166–168 In one system of particular note, the NdN bond can be broken photochemically, and then oxidation of the nitride product regenerates the NdN bond.169 These exemplify the ability to control the reactivity of N2 using light and metal oxidation state.

MLn

MLn

MLn

N

N

N

N

N

N

MLn

MLn

MLn

Fig. 26 Cleavage of the NdN bond of N2 to terminal nitride complexes through a nonlinear “zigzag” transition state.

1.17.3.5.2

N2 splitting to bridging nitrides

Splitting of N2 to give bridging nitrides is relevant to N2 activation that occurs at zerovalent metal surfaces in the Haber-Bosch process, and therefore homogeneous analogues have been studied. The isolated products may be bimetallic, trimetallic or tetrametallic, and have a wide variation in geometry and oxidation state. In these systems, it is evident that the energetic price of NdN cleavage is repaid through the formation of many M–N bonds, and typically by the oxidation of the metals from an unusually

tBu

(tBuO)2Si

O

O

K

O

O

tBu

tBu

(tBuO)2Si

Si(OtBu)2

t O U N U O Si(O Bu)2 tBu O O K O O tBu t K ( BuO)2Si Si(OtBu)2 O O tBu tBu

R iPrN

Cp*

M

Solid state

Cp* N N

M iPrN

NiPr

iPrN

NiPr R

K

O

O

tBu

Si(OtBu)2

(tBuO)2Si

R

NiPr

O

N t U N U O O Si(O Bu)2 N tBu O O K O O tBu (tBuO)2Si Si(OtBu)2 K O O tBu tBu

N2

(tBuO)2Si

O

Cp*

M

N

N

Cp* M iPrN

NiPr R

Fig. 27 Examples of N2 cleavage to give bridging nitride complexes. M ¼ V, Nb, Ta; R ¼ Me, Ph.

low oxidation state. The transition metal systems that accomplish this are most typically early transition metals or actinides, and two representative examples are shown in Fig. 27.170,171 In some cases the low-valent precursors are not isolated. One example with iron was shown above in Fig. 22. In an additional example shown in Fig. 28, KH is added to reduce the vanadium, and N2 is subsequently bound and cleaved without observation of an intermediate.172 More systems have been seen where hydrides are reductively eliminated to enable N2 binding, which is advantageous because of the milder reducing agent.133,160,173,174

536

Dinitrogen Binding and Functionalization

tBu

O

tBu

K

O tBu

DME

O

V

O

tBu

NTol

4 KH, THF, N2 – H2

O

Tol N V

tBu

V

O

tBu

O

K

tBu

tBu

N Tol

N

K

DME

DME

K

N

DME

tBu

Fig. 28 Example of KH as a reducing agent that enables N2 binding and cleavage. DME ¼ 1,2-dimethoxyethane.

1.17.4

Functionalization of N2

1.17.4.1

Forming NdH bonds at coordinated N2

1.17.4.1.1

Protonation of N2

In 1972, Chatt and coworkers demonstrated that terminal N2 complexes can be protonated using strong acids.175 For instance, trans(dppe)2W(N2)2 (dppe ¼ Ph2PCH2CH2PPh2) and trans-(dppe)2Mo(N2)2 can be protonated using strong inorganic acids (e.g. HBr or HCl) to give products that were first suspected to be the corresponding diazene complex MX2(dppe)2NH]NH based on NMR spectroscopy, but later turned out to be hydrazido(2–) complexes from double protonation of the distal nitrogen (Fig. 29).176,177 It was also possible to observe a diazenido complex from deprotonation of the hydrazido(2–) species.178 Analogous formation of hydrazido(2–) species was observed for Mo and W complexes bearing monodentate phosphines.23 There were also differences based on the stereochemistry: trans-W(N2)2(MePPh2)4 with HCl led to partial N2 protonation like that in Fig. 29, but the cis isomer quantitatively formed NH3.177,179 A mixture of NH3 and N2H4 was observed from the cis-Mo(N2)2(Me2PPh)4 analogue.180 These studies demonstrated the higher reactivity of cis isomers toward NH3 formation. The higher reactivity of monodentate phosphine complexes with respect to their bidentate analogues was attributed to the easier dissociation of a monodentate PMe2Ph ligand upon protonation of N2.179 This principle was used to achieve protonation under milder conditions (HBr in toluene) of cis-[Mo(N2)(PPP)(PPh3)] (PPP ¼ PhP(CH2CH2PPh2)2) (Fig. 30).181,182 This complex produces NH3 (0.7 equiv./Mo) and [Mo(PPP)Br3] (>90% yield), indicating a stoichiometry in which the Mo center provided three electrons. The facile dissociation of the monodentate PPh3 may provide a vacant site for binding of the anionic Br− to further increase the electron density at the metal center and enable N2 protonation. Similar reactivity in lower yield was also found for low valent FedN2 complexes (Fig. 31). Leigh and coworkers reported the stoichiometric reduction of N2 to NH3 using Fe(dmpe)2N2 (dmpe ¼ Me2PCH2CH2PMe2).183–185 After treating the complex with HCl in THF, sub-stoichiometric amounts of NH3 (0.12 equiv) and traces of N2H4 were detected. Similarly limited NH3 producing

N Ph2 N P M P Ph2 N N

Ph2 P P Ph2

HX

NH Ph2 N Ph2 P P M P P Ph2 X Ph2

Ph2 N P M P Ph2 X

HX base

Fig. 29 Protonation of N2 at bis(diphosphine) complexes. M ¼ Mo, W; X− ¼ HSO−4 , Br−, Cl−.

Ph

P

PPh3 PPh2 Mo N N

PPh2N N Fig. 30 Protonation of cis-Mo(N2)(PPP)(PPh3) to produce NH3.

Br HBr Toluene

+

NH2

NH3 +

Ph

P

Mo

PPh2Br

PPh2 Br

Ph2 P P Ph2

X–

Dinitrogen Binding and Functionalization

NH3

537

+ H2N NH2

TfOH Et2O

NH3

PR2

R2P

HCl THF

TfOH Pentane

Fe N N

R2P

H2N NH2

PR2

R = Ph, Me, (CH2)3OMe Fig. 31 Protonation of N2 at iron bis(diphosphine) complexes.

ability was observed with other iron-diphosphine complexes such as Fe(dppe)2N2 and the water soluble Fe(DMeOPrPE)2N2 (DMeOPrPE is the ligand where R ¼ (CH2)3OMe).185,186 The overall six-electron reduction to NH3 was proposed to occur via the participation of multiple Fe0 complexes acting as sacrificial reducing agents to provide the six electrons, as further suggested by the isolation of the bridging N2 dimer [Fe(dmpe)2]2(m-N2).187 Fe(dmpe)2N2 reacted with triflic acid (TfOH) to give N2H4 in pentane, NH3 in THF, and a mixture of N2H4 and NH3 in Et2O, and the solvent dependence was attributed to the modulation of the acidity of TfOH by the different solvents.187 Protonation of complexes with bridging end-on/end-on N2 has also been investigated. In a particularly insightful study, Bercaw and coworkers reported the protonation of partially 15N-labeled [Cp 2Zr(N2)]2(m2-N2) with HCl, which forms N2H4 (Fig. 32).125,188,189 They showed that the amount of 15N label in the N2H4 could be rationalized through protonation of the terminal N2 ligands that is more rapid than the bridging N2 ligands. It has since emerged that this trend is general: terminal N2 is more reactive than end-on/end-on bridging N2, not only for protonation190 but also for N2 silylation (see Section 1.17.4.2.1).191 Lewis acids, in addition to their ability to cause ground-state weakening (Fig. 9 above), can enhance the reactivity of M–N2 complexes. This was recently demonstrated by Szymczak and coworkers in the reaction of [(depe)2Fe(N2)] (depe ¼ Et2PCH2 CH2PEt2) with acid, which only gives N2 protonation if done in the presence of B(C6F5)3: in this case it was possible to isolate the hydrazido(2−) complex [(depe)2Fe(NNHBAr3)]+ (Fig. 33).96 Similar observations were made by Etienne and coworkers with group 6 (dppe)2M(N2) complexes (M ¼ W and Mo).192

1.17.4.1.2

NdH bonds from H2

The examples above used protons from a Brønsted acid and electrons from the M–N2 complex. However, protons and electrons can instead be supplied using H2 or hydrides, which provide one or two electrons along with the proton. Organometallic chemists have

N N Zr N N Zr N N

HCl Toluene -80 °C

H2N NH2

Fig. 32 Protonation of a dizirconium N2 complex with bridging and terminal N2 ligands.

F

F Et2P Et2P

PEt2 Fe N N PEt2

B(C6F5)3

Et2P Et2P

PEt2 Fe N N F

PEt2 F

F

F

F

F B FF F F

HBArF4

Et2P Et2P

F F

F

PEt2

F

F

HF F Fe N N B F F PEt2 F F F F F F F F F

Fig. 33 Lewis acid mediated activation of N2 on an iron bis(diphosphine) complex. ArF ¼ 3,5-bis(trifluoromethyl)phenyl.

BArF4

538

Dinitrogen Binding and Functionalization

noted that this route has advantages because it might bypass the need for the usual strong acids and strong reducing agents.7,174,193 However, this route provides less driving force, as the net hydrogenation of N2 is only slightly exothermic and is entropically unfavorable, and thus the demands on the catalyst are greater. H2 is a compelling proton/electron source, but it is relatively unreactive and can also displace N2. However, H2 complexes like [CpRu(Z2-H2)(dtfpe)]BF4 (dtfpe ¼ (p-CF3Ph)2PCH2CH2P(p-CF3Ph)2) and in situ generated trans-[(dppe)2RuCl(Z2-H2)]X are capable of hydrogenating the coordinated N2 molecule in trans-(dppe)2W(N2)2, to produce the hydrazido(2–) complex [trans(dppe)2WF(NNH2)]BF4 and subsequently NH3 (Fig. 34).194,195 Mechanistically, this reaction bears similarity to the attack of a Brønsted acid on the tungsten-N2 complex described above, except that the acid is the dihydrogen complex. H2 acts as the source of the proton, but in this case the electrons still are supplied by the metal center bound to N2. The first examples of direct hydrogenation of N2 complexes by H2 used Zr.11 The treatment of a side-on bridging N2 zirconium complex (P2N2)Zr(m-N2)Zr(P2N2) (P2N2 ¼ [(PhPCH2SiMe2)2N]2) with H2 generated [(P2N2)Zr]2(m-Z2-N2H)(m-H), which contains singly protonated N2 (Fig. 35, top).196,197 In a related reaction, an N2-bridged zirconocene complex (CpMe4H)2Zr (m-N2)Zr(CpMe4H)2 reacts with 1 atm of H2 at room temperature to give (CpMe4H)2Zr(m-N2H2)Zr(CpMe4H)2 (Fig. 35, bottom).121 Further heating under H2 led to the cleavage of the NdN bond, forming NH3. Overall, these N2 complexes are able to react with H2, producing new NdH bonds. Alternatively, hydrogenation of N2 can be carried out using metal hydride complexes that are pre-formed from H2. For example, (PNP)Ti(CH2SiMe3)2 (PNP ¼ N(4-Me-2-PiPr2-C6H3)2) and H2 gives [(PNP)Ti]2(m-H)4, which readily coordinates N2 with release of one equivalent of H2 and forms the side-on/end-on N2 bridging complex [(PNP)Ti]2(m-Z1:Z2-N2)(m-H)2 (Fig. 36).133 Upon heating, [(PNP)Ti]2(m-Z1:Z2-N2)(m-H)2 further transforms to [(PNP)Ti]2(m-NH)(m-N)H. Similar reactivities were also found for other titanium polyhydride complexes as well as from chromium hydride complexes.124,173 Other transition metal hydrides such as Cp2ZrHCl, H2Fe(CO)4,198 or HCo(CO)4199 react with cis-W(N2)2(PPh2Me)4 to produce stoichiometric NH3 and N2H4 (Fig. 37). The mechanisms are not clear, and with the acidic hydrides the initial step may be protonation by the hydride.

N Ph2 N P W P Ph2 N N

Ph2 P P Ph2

R2 P

+ 2

P R2

BF4 Ru H H [Ru]H2

R2 P P R2

H

N Ph2 N P W P Ph 2 F

BF4 Ru

H H

THF 25 °C, 1 h –2 [Ru]H

BF4

H

BF3.THF excess [Ru]H2

Ph2 P P Ph2

Fig. 34 N2 hydrogenation mediated by a ruthenium phosphine complex using molecular hydrogen.

Me2 Me2 Si Si

Me2 Me2 Si Si

Ph

P

Si N N Si Me2 Me2 P Zr

P

H2

N

N Ph

Ph

Ph

Si N N Si Me2P P Me2 Zr H N N

-80 °C Ph

Ph

Zr

P

H P

Zr

Zr

Ph P

Si Si N N Me2 Me2 Si Si Me2 Me2

Si Si N N Me2 Me2 Si Si Me2 Me2

N Zr N

H2 22 °C

Fig. 35 Hydrogenation of side-on/side-on zirconium-N2 complexes with H2.

Ph

H N Zr N H

Zr

H2 80 °C

NH3

NH3

Dinitrogen Binding and Functionalization

[Ti] [Ti] =

CH2SiMe3

H2/N2

CH2SiMe3

60 °C H2/N2

H2 iPr iPr

N P

i P Pr iPr

Ti

H H H H

[Ti]

H [Ti]

N

[Ti]

60 °C

RT N2

[Ti]

H N

539

[Ti]

-H2

N N H H

[Ti]

Fig. 36 N2 hydrogenation using titanium polyhydride complexes derived from H2.

PPhMe2 PPhMe2 W PPhMe2 N N

Me2PhP N N

[M-H] =

Zr

[M-H]

H

OC

Cl

OC

NH3 + H2N NH2

H

H

Fe

CO

CO

OC

H Co

OC

CO CO

Fig. 37 N2 hydrogenation on tungsten complexes by metal hydrides (illustrated at bottom).

1.17.4.1.3

Catalytic N2 reduction to NH3

The catalytic reduction of N2 to NH3 will be discussed only briefly, since a recent comprehensive review is available for the interested reader.32 In homogeneous systems, three pathways are most often discussed for N2 reduction to NH3, and these are outlined in Fig. 38. The steps are labeled with H+/e− because these overall pathways do not specify the order of proton and electron transfer (or whether they are concerted). Two of the pathways are initiated by N2 binding and H transfer to the distal N atom. The alternating pathway (purple) involves sequential NdH bond formation at the distal and proximal N atoms in turn. Formation of N2H4 is often thought to imply the alternating pathway, since this is the only one having coordinated N2H4 as an intermediate. The distal pathway

M

+

M N N

N N

+M

H+/e–

Alternating

H M N NH

N2 Cleavage

H+/e–

H+/e–

M N NH

Distal

M N NH2

M N NH2 H

M N N M

H+/e–

–M N H+/e– -NH3

H+/e–

H+/e–

M N NH2 H2

H+/e– -NH3

M NH2 H+/e–

M Fig. 38 Pathways for formation of NH3 from N2.165

M N

+

NH3

H+/e–

M NH

540

Dinitrogen Binding and Functionalization

(blue) involves repeated NdH bond formation at the distal N atom until NdN cleavage yields NH3 and leaves behind a metal nitride (this is most often drawn as a terminal nitride, but it could be bridging in multimetallic pathways). Subsequent H transfer to the nitride gives NH3, and these steps are not discussed here as excellent reviews are available on nitride reactivity.163,200,201 The third pathway involves cleavage of the NdN bond prior to H transfer, and is termed the N2 cleavage mechanism (or dissociative mechanism) and is colored red in Fig. 38.165 This pathway bypasses the high-energy MNNH intermediate, but requires an extremely reducing pair of metals that can accomplish the net six-electron reduction of N2 in one step (three electrons per metal). Molybdenum and iron systems discussed in Section 1.17.3.5 above can engage in N2 splitting to form terminal and bridging nitrides, respectively. A particular challenge for catalysis via this pathway is that the splitting is often thermodynamically driven, such that the nitride product is extremely stable and inert, and thus does not continue to the further steps needed to complete a catalytic cycle .165 However, this challenge has been overcome in a few systems: for example, an N2-derived rhenium nitride complex could produce ammonia using the reductant/acid combination SmI2/H2O,168 which provides high-energy coupled protons and electrons as described in the following section.

1.17.4.1.4

NdH bond strengths

Quantitative analysis of the thermodynamics of forming NdH bonds from metal–N2 species involves consideration of the NdH bond dissociation free energies (BDFE) in diazenido, hydrazido, ammine, amido, and imido complexes. The BDFEN–H values of some metal–NxHy complexes (mostly from computations) are given in Table 3. Experimentally, NdH BDFEs are usually determined from pKa and E0 values using the relation BDFEeff (in kcal mol−1) ¼ 1.37pKa + 23.06E0 + CG (where CG is a constant that is dependent on the solvent).218 Typically, pKa values are determined by acid-base titrations against standards of known reference values, while E0 are determined from cyclic voltammetry. The reported BDFEN-H for MNNHy complexes are all under 50 kcal mol−1 and are typically lower than the reported BDFEN-H for M–NHy. In accordance with their highly reactive nature, very few M-NNH have been reliably characterized.219,220 In the context of NH3 production, the only experimental BDFEN–H for a M–NNH species was determined in (HIPTN3N)MoNNH (HITPN3N ¼ (3,5(2,4,6-iPr3C6H2)2C6H3)NCH2CH2)3N, Fig. 39).214,219 The cyclic voltammogram of (HIPTN3N)MoN2 shows a Mo3+/2+ couple at −1.81 V (vs. Fc+/Fc) in THF, and the pKa of (HIPTN3N)Mo–NNH was estimated to be close to the pKa of DBU–H+ (18.5).221,222 These give a BDFEN–H of 47.0 kcal mol−1 for (HIPTN3N)MoNNH (using CG in THF ¼ 59.9).223 This value agrees with Table 3

Selected NdH bond dissociation free energies of NxHy ligands.a

Complexb

Metal-NNHy [(HIPTN3N)Mo(NNH) BP3Fe(NNH2) SiP3Fe(NN(Me)H) trans-(dppe)2IMo(NN(Cy)H) trans-(dppe)2(MeCN)Mo(NN(Cy)H) trans-(dppe)2(3,5-(CF3)2C6H3CN)Mo(NN(Cy)H) Metal-NHy cis-(PONOP)Re(NH2)Cl2 (PNP)Ir(NH2) trans-[(Ph-tpy)(PPh2Me)2Mo(NH3)]+ tBu + cis-[Cp(PPh 2 N2 )Mo(NH3)(CO)] [(PY5)Mo(NH3)]2+ Ph + [Cp (PtBu 2 N2 )Ru(NH3)] [(tpy)(NMe2bpy)Ru(NH3)]2+ (TMP)Ru(NH3)2 [(tpy)(NMe2bpy)Fe(NH3)]2+ [(PhNCH2CH2)3N]Mo(NH3) [(BP3)Fe(NH2)]+ (salen)Mn(NH3) (F)(H2PCH2CH2PH2)2Mo(NH3) (
5-C5Me4SiMe3)2Ti(NH3) a

Solvent

BDFENdH (kcal mol−1)

References

NH3

NH2

NH

THF THF THF THF THF THF

– – – – – –

– 47 – – – –

43 (47)b 35 48 33 35 36

202 203 204 205 205 205

THF gas phase THF Et2O MeCN THF THF C6H6 THF – Et2O gas phase benzene gas phase

– – 46 84 68 83 79 82 82 52 – 85 41 42

78 95c 64 61 65 89 86 93 90 64 80 84 92 79

43 71c – 73d 64 72 – 75 – 42 65 60 37 –

168 206 207,208 209 210 211 212 213 212 214 161 215 216 217

From computations unless otherwise noted. For specifics of the supporting ligands, see the cited publications. Abbreviations: HIPTN3N: (3,5-(2,4,6-iPr3C6H2)2C6H3)NCH2CH2)3N; BP3: (B(o-C6H4PiPr2)3)−; SiP3: (Si(o-C6H4PiPr2)3)−; dppe: Ph2PCH2CH2PPh2; PONOP: 2,6-bis(diisopropylphosphinito)pyridine; PNP ¼ N(CHCHPtBu2)2; PY5: 2,6-bis(1,1-di(pyridin-2-yl)ethyl)pyridine; tpy: terpyridine; TMP: tetramesitylporphyrin; bpy: bipyridine; salen: N,N0 -bis(2,4-di(t-butyl)salicylidene)-trans-1,2-diaminocyclohexane. c Measured bond enthalpies. d Measured after loss of CO ligand following initial NdH bond oxidation. b

Dinitrogen Binding and Functionalization

HIPT HIPT N N

N N Mo

E0 = –1.81 V + e–

HIPT N

HIPT N N

– e–

N

N N

HIPT

541

– HIPT

Mo

N

N

+ H+ – H+ pKa ~ 16.6 BDFEN-H ~ 47 kcal mol–1

N N

HIPT HIPT N N

Mo

H HIPT N

N HIPT HIPT N N

H2

– N2

H Mo

HIPT N

N

Fig. 39 Using the measured redox potential and acidity to calculate the BDFEN–H of (HIPTN3N)MoNNH (HIPT ¼ 3,5-(2,4,6-iPr3C6H2)2C6H3), also showing the decomposition to release H2.

computational studies.203 It is important to note that below 51 kcal mol−1, NdH bonds are thermodynamically prone to bimolecular release of H2 (BDFEH–H ¼ 102 kcal mol−1 in MeCN).218 This is one reason why production of H2 from MNNH species during N2 reduction is a major challenge.

1.17.4.1.5

Concerted PCET

The addition of protons or electrons builds up charge, and it is possible to avoid this energy cost by adding protons and electrons simultaneously. This strategy of “proton-coupled electron transfer,” abbreviated PCET, often improves reaction rates.224 PCET can be stepwise (PT-ET or ET-PT) or concerted (occurring in one elementary step), and these are often difficult to distinguish experimentally, especially in catalytically competent homogeneous systems under discussion here.32,218 For instance, it had been proposed that the (HIPTN3N)ModN]NH complex involved in the catalytic triamidoamine-molybdenum system is formed via stepwise proton-transfer/electron transfer (PT-ET) pathway, but Peters and coworkers showed that concerted PCET is a more feasible route to (HIPTN3N)ModN]NH (Fig. 40).21,32

HIPT HIPT N N

N N Mo

HIPT N

N LutH+ (PT)

HIPT H (concerted PCET)

HIPT N N [LutH]+ / CrCp*2 or CoCp2

HIPT HIPT N N

N N Mo

H

HIPT N

N Fig. 40 Mechanistic alternatives for the formation of (HIPTN3N)ModN]NH during catalysis.

N N Mo

H HIPT N

N

reductant (ET)

542

Dinitrogen Binding and Functionalization

R

R e–

N H

R

R R = H, Me

R BDFEN-H ~ 35 kcal/mol H

CoII

N H

R

H+

CoII

H

+

+ e–

BDFEC-H < 29 kcal/mol Fig. 41

CoI 'G(H–) < 41 kcal/mol

Top: ET of pyridinium acids and BDFEN–H of the corresponding pyridinyl radicals. Bottom: PT-ET for decamethylcobaltocene and corresponding energetics.

In a related opportunity for PCET, the combination of strong pyridinium acids and metallocene reductants can lead to the in situ formation of pyridinyl radical species with weak NdH bonds (BDFEN–H  35 kcal mol−1, Fig. 41). These pyridinyl radicals could therefore use PCET to deliver energetic hydrogen atom equivalents to the nitrogen-containing ligands during the catalytic cycle.214 Using the same combination of acids and reductants, PCET was proposed in Peters’ (BP3)Fe catalytic system.32 Based on computational and experimental evidence, they surmised that even with pyridinium acids, reduction of the pyridinium rings is less favorable than proton transfer from the pyridinium to the metallocene reductant. The relevant metallocenes (Cp2Co and Cp 2Co) are predicted by DFT to undergo ring protonation to generate species that have BDEC−H  30 kcal mol−1), suggesting they are competent PCET donors.203,225 Given the particular effectiveness of metallocenes in N2 fixation catalysis, metallocene-mediated PCET should be considered a likely pathway for N − H bond formation.32 Another strong PCET donor for N2 activation can be formed from the strong reductant SmI2 and H2O (or alcohols). Accordingly, Nishibayashi and coworkers have used SmI2/ethylene glycol for catalytic N2 reduction with high selectivity (greater than 90% for NH3) from a Mo catalyst.226

1.17.4.2

Forming other bonds at coordinated N2

Ammonia is not the only nitrogen product that is desired from N2! Formation of N–X bonds, where X represents other nonmetals, is a potential route directly from air to functionalized N-containing products without the intermediacy of ammonia or hydrazine. Functionalization of N2 via transition metal nitrides (derived from N2 splitting; see Section 1.17.3.5 above) is a similarly rich and extensive research area; readers may refer to reviews on the topic.200,201

1.17.4.2.1

NdSi bonds

As described above, isolation of M–NxHy intermediates is often challenging due to kinetically facile proton or H• transfer, which is driven by the weak NdH bonds in these intermediates as described above. In order to stabilize analogues of these intermediates, chemists have leveraged cationic trialkylsilyl groups as proton surrogates.220,227–230 Thus, the formation of NdSi bonds from N2 has been an active area of research.18,190,231 The most common strategy for forming NdSi bonds to N2 involves adding a silyl halide or pseudohalide (R3SiX) to a metal-N2 complex. End-on/end-on bridged N2 complexes are generally not reactive, and a recent direct comparison of b-diketiminate-supported iron(0) N2 complexes shows that they react only when they can convert to terminal end-on N2 species.191 In terminal N2 complexes, the enhanced nucleophilicity of the distal N atom leads to NdSi bond formation to give either a monosilylated product (silyldiazenido) or a disilylated product (disilylhydrazido(2–)), with the first examples from Hidai and coworkers in the 1980’s.14 Some more recent examples are shown in Fig. 42. For example, silylation of the molybdenum N2 complex [Na(15-crown-5)][Cp Mo(depf )(N2)] (depf ¼ Fe(C5H4PEt2)2) with 1 equiv. Me3SiCl gives single addition in the silyldiazenido complex Cp Mo(depf )NNSiMe3.232 In contrast, the addition of 1 equiv. of Me3SiCl to a chromium bis(dinitrogen) complex gives two NdSi bonds in a disilylhydrazido(2–) complex.233 Increasing the steric bulk of the R group on the silyl (pseudo)halide should also disfavor double silylation,234 although a systematic study of this effect has not yet been reported. One way to encourage double silylation is to use a bis(silyl)ethane reagent, because the second silyl addition becomes intramolecular.235,236 In an instructive recent example, reduction of the iron bromide complex (BP3)FeBr (BP3 ¼ B(o-iPr2PC6H4)3) in the presence of excess Me3SiCl under an N2 atmosphere gave the monosilylated diazenido complex [Na(THF)][(BP3)FeNNSiMe3] (Fig. 42, bottom left). Replacing Me3SiCl with 1,2-bis(chlorodimethylsilyl)ethane gives instead the disilylhydrazido(2–) complex (Fig. 42, bottom right).

Dinitrogen Binding and Functionalization

Et2 P Fe

Cp* Mo

PEt2

N

N

Fe

O

R2P R2P

N

Na

Fe

THF

N

N SiMe3

Et [Cr] Cl

Et Et

Cy P Cr N Cy

1 atm N2 xs Na/Hg xs Me3SiCl

hydrazido(2–) SiMe3

N SiMe3

1 atm N2 xs Na/Hg

Br R2P R2P

diazenido

Et

+

Me3Si N

Cp* Mo

PEt2

O O Na O O

K(2.2.2-crypt)+ – Et Et 1 equiv Et Me3SiCl Et per Cr Cy P Cr N2 Cy N2

2

Et2 P

1 equiv Me3SiCl

PR2

Fe

Me2Si SiMe2 Cl Cl

Si R2P R2P

Si

N N

PR2

Fe

B

B

543

B

R2P R = iPr Fig. 42 Examples of N2 silylation using silyl electrophiles.

One complication from using Me3SiX in the presence of a strong reducing agent is the potential for these to react directly to form silyl radicals. Therefore, there is often ambiguity about whether the NdSi bond formation occurs as reduction of the N2 complex followed by a nucleophilic attack at Si or whether the reductant generates silyl radicals that subsequently attack N2. Nishibayashi and coworkers used DFT to calculate the feasibility of various pathways for generation of reactive silyl species from Me3SiCl.237 These calculations indicated that direct cleavage of the SidCl bond is not feasible, but reduction of Me3SiCl to give Me3SiCl− followed by cleavage to Me3Si● and Cl− was calculated to be exergonic by 20.5 kcal mol−1. This suggests that attack of N2 by silyl radicals formed via reduction of Me3SiX is the most likely pathway. There are also examples of N2 complexes reacting with silanes (R3SiH) to form new NdSi bonds, often with concomitant formation of a metal hydride. For example, Semproni et al. reported the reaction of a side-on bridged dihafnocene N2 complex with CySiH3 to yield an equimolar mixture of two isomeric silyldiazenido complexes (Fig. 43, top).238 Interestingly, heating a solution of these isomers led to cleavage of the NdN bond of the diazenido to give a [NSi(H)CyNH]3− fragment, accompanied by the formation of a terminal HfdH bond.

Cy

Hf

N N

Hf

CySiH3

SiH2 H

N

N

Hf

Br R2P R2 P

Fe B

Me2Si SiMe2 Cl Cl

Hf

H

1 atm N2 3 equiv Na/Hg

N

75 °C, 18 h

Hf

SiHCy N H

Si

N

Si Si N

Si

N R2P R2 P

Fig. 43 Formation of new NdSi bonds from coordinated N2 using silanes.

Hf

Fe B

PhSiH3

R2P R2P

N SiH2Ph

Fe H B

544

Dinitrogen Binding and Functionalization

Reactivity with a silyl halide and a silane in series to form NdSi bonds is exemplified by research from Peters and coworkers using an iron complex supported by a bis(phosphino)borane ligand (Fig. 43, bottom).239 Reduction of the iron(I) bromide complex (BP2)FeBr (BP2 ¼ PhB(o-iPr2PC6H4)2) with Na/Hg amalgam in the presence of 1,2-bis(chlorodimethylsilyl)ethane under an N2 atmosphere yielded the corresponding disilylhydrazido(2–) complex. This complex could react further with PhSiH3 to form a new NdSi bond through formal hydrosilylation of the FedN bond. This trisilylhydrazido complex could also be accessed in a one-pot method starting from (BP2)FeBr. Finally, we move to catalytic reactions. Chemists have described a number of cases where the addition of excess reductant and silyl halide gives catalytic formation of trisilylamines. The Nishibayashi group has been most active in this area, demonstrating catalytic silylation of N2 to N(SiMe3)3 with a variety of transition metal complexes supported by pincer ligands,240 including Fe,241,242 Co,242 Mo,243 Rh,244 Re,245 and Ir (Fig. 44).246 Interestingly, they found that catalytic N2 silylation can be mediated by simple, coordinatively saturated organometallic complexes such as binary transition metal carbonyl complexes and metallocenes.247,248 These simple complexes were shown to be just as catalytically competent as some of the pincer complexes, which indicates that the actual catalyst may be generated by loss of supporting ligand(s). The catalytic competence of these simple complexes complicates the interpretation of catalytic N2 silylation using other complexes, since impurities may be catalysts as well. Another notable example of catalytic silylation of N2 was reported by Lu and coworkers, in which a dicobalt system mediated the reaction using Me3SiCl as a silylating reagent and KC8 as a reductant.249 Using a catalyst loading of 0.05 mol%, addition of Me3SiCl and KC8 (2000 equiv. each) in THF at room temperature for 12 h gave a 30% yield of N(SiMe3)3, with a TON of 195. When the reaction mixture from one run was filtered to remove any heterogeneous decomposition products and subsequently resubjected to catalytic conditions, an even higher TON of 316 was achieved, suggesting that the catalyst is soluble. Though only the dicobalt starting material and the initial N2 complex could be isolated, DFT calculations were used to propose the overall catalytic cycle as shown in Fig. 45. Like Nishibayashi, they proposed that the silyl group adds as a silyl radical, an idea that was supported by the formation of Me3SiSiMe3. However, clearly more research is needed to have strong experimental support for the mechanism(s) of trisilylamine formation in the growing number of catalytic systems.

N2 (1 atm)

6 Na (600 equiv)

catalyst

6 Me3SiCl (600 equiv)

THF r.t., 20 h

2 N(SiMe3)3

6 NaCl

Fe(CO)5

R

Si H

PCy2 PMe3 Fe N N PCy2

15 (R = Me) 26 (R = Ph)

PtBu2 N Fe L

25, 2a Fe2(CO)9

PCy2 R

Si Co

9 PtBu2 33 (L = N2) 14 (L = Cl) 4 (L = H) 3 (L = Me)

Fe3(CO)12 12 FeCp2 13

PCy2

N N PR3

Co2(CO)8 22a, 36a,b

CoCp2 28 (PR3 = PMe3) 23 (PR3 = PMe2Ph) 31, 41b (PR3 = PMePh2)

8a,b

Fig. 44 Catalytic silylation of N2 by Fe and Co complexes; the numbers indicate the TON of N(SiMe3)3 per metal atom. a ¼ in 1,2-dimethoxyethane rather than THF. b ¼ 40 h rather than 20 h.

Dinitrogen Binding and Functionalization

545

Fig. 45 Computationally proposed catalytic cycle of N2 reduction to N(SiMe3)3 by a dicobalt system. Me3Si● is thought to be formed from the reduction of Me3SiCl by KC8.

1.17.4.2.2

NdC bonds

The formation of new NdC bonds from N2 is an avenue to organic N-containing compounds without the intermediacy of NH3. Catalytic NdC bond formation from N2 has yet to be realized, although several systems have shown promising results that have been reviewed.18,231,250 One of the most common methods of stoichiometric NdC bond formation from N2 is to use carbon electrophiles, such as alkyl and acyl halides, which can react with the nucleophilic distal nitrogen atom of a terminal M–N2 complex. Cummins and coworkers demonstrated this reactivity through the addition of MeOTs and benzoyl chloride to molybdenum tris(amido) N2 complexes (Fig. 46, top).144 Further, a report by Xi and Zhang demonstrated the use of both of these types of electrophiles in one pot (Fig. 46, bottom).115 Addition of potassium and MeOTf to the side-on bridged discandium N2 complex gave a dimethylhydrazido(2–) complex, and further addition of acyl chlorides or alkyl halides gave the corresponding free tetrasubstituted hydrazine with formation of the dimeric scandium halide complex. Attack of ligated N2 on aryl electrophiles is rare, with only a handful of examples (Fig. 47). In the first, reaction of a bis(dinitrogen)molybdenum complex ligated by a crown thioether with various aryl halides led to the isolation of the respective aryldiazenido complexes.251 Hidai and coworkers also reported bimetallic N2 arylation using coordinated aryl halides, in which coordination to a transition metal engendered greater electrophilicity of the aryl halide.252,253

546

Dinitrogen Binding and Functionalization

Ph

Na(THF)x

Me N

N

MeOTs

N

R R R N Mo N Ar N Ar Ar

O Ph

N

R R R N Mo N Ar N Ar Ar

R = tBu

N

Cl

N

R R R N Mo N Ar N Ar Ar

R = Ad

Ar = 3,5-dimethylphenyl

K

nBu

2 KX

Sc

N NiPr

iPrN

N2 + 3 K

N

iPrN

Sc

2 MeOTf + K

NiPr

2 KOTf

nBu

Me

nBu

iPrN

Sc

X

iPrN

X NiPr

Me N

NiPr

Sc

iPrN

N

Sc NiPr

Sc

EX

Me

Me N

N

E

E

NiPr

iPrN

nBu

nBu

nBu

O

EX = BnBr, Ph

O Cl

Cl , Cl O

Fig. 46 NdC bond formation from N2 using alkyl and acyl electrophiles. Ad ¼ 1-adamantyl.

N S S

Mo

N

X

N S

N

S

S

X = Br, I

Mo

S

N

N Ph2 N P W P Ph2 N C S Fig. 47 Reactions of bound N2 with aryl halides.

Bu4N+ – Ph2 P P Ph2

S S

X

N

CO OC Co OC F

CO2Me

OC Co CO OC

N Ph2 N P W P Ph2 N C S

O

Ph2 P P Ph2

CO2Me

Dinitrogen Binding and Functionalization

547

Because bound N2 is not electrophilic, there are no well-characterized examples of reactions with nucleophiles. This differs from metal-carbonyl complexes, where migration of alkyl groups is commonplace. To our knowledge, the only report claiming attack of N2 on a carbon nucleophiles has only partial characterization.254–256 The manganese piano stool complex CpMn(CO)2N2 was reported to react with MeLi or PhLi to give unstable species that were assigned as the respective diazenido complexes. Further reaction of these species with H2SO4 or Me3OBF4, respectively, gave neutral diazene complexes which were also partially characterized. In another class of reactions, coordinated N2 has been observed to react with (hetero)allenes. For example, a dinuclear tris(aryloxide)methyl titanium N2 complex reacts with CO2 and isocyanates (Fig. 48).257 Addition of CO2 yielded a trisubstituted hydrazido complex via addition of three molecules of CO2. On the other hand, only two molecules of isocyanate add to N2, giving the 1,2-disubstituted hydrazido(2–) complex. Phenylallene gives only one addition (proposed to be formed via [2+ 2] cycloaddition of the C]C bond to the M]N bond), and two isomers of the protonated product are isolated.

K4(DME)34+ 4– C O O O Ti O

O

O C Ti OO O

N

H3C O

O

N

CO

2

O

(1

K(THF)6+

O tBu

t BuN

CO

tBu

tBu

C Ti OO

–

ba r)

O

tBu tBu

O

O

tBu tBu

K

K

O

Ti N N Ti

O THF tBu

Ph tBu

O

O THF tBu

K

tBu

xs •

K(THF)6+ THF K K THF

N N O

Ti O O C

H –

Ph

tBu

O

THF THF tBu O K O N THF OO C Ti N K K N Ti C OO THF N O O K tBu THF

–

Ph

N

C Ti OO

N O

K(THF)5+

CH2 K

THF K THF

Ti O O C

Fig. 48 Addition of allene, isocyanate, and CO2 to a dititanium N2 complex.

Products from cycloaddition are more commonly observed with side-on bridged N2 complexes.231 For example, alkynes add to the side-on bridged dizirconium N2 complex to give the vinylhydrazido(3–) complex shown on the right of Fig. 49.258 This reaction is proposed to proceed via the cycloaddition with one alkyne to give an unobserved zirconaheterocyclobutene intermediate, which subsequently reacts with another equivalent of alkyne to yield the isolated product.

Me2 Me2 Si Si

Me2 Me2 Si Si

Ph

Si N N Si P P Me2 Zr Me2

Ph

P

Zr

[Zr]

Ph Ar

N

N

Ar

P

Ar = Ph, p-Tol, Ph p-tBuC6H4

N

N

[Zr]

Si N N Si Me2 Me2 Si Si Me2 Me2 Fig. 49 Alkyne addition across a ZrdN bond with postulated intermediate.

C

Ph C

H

Ar Ar Ph

Si N N Si P P Me2 Zr Me2 Ph H N Ar N C H P

Zr

P

Si N N Si Me2 Me2 Si Si Me2 Me2

Ph

548

Dinitrogen Binding and Functionalization

The formation of new NdC bonds via insertion of CO into the NdN bond of N2 has also been explored (Fig. 50). In a report from the Chirik group, reaction of a side-on bound N2 dihafnium complex with CO yielded either a bridging oxamido complex or an imido cyanate complex depending on the amount of CO added.259 Further reactivity of related N2 and CO-derived metallocene nitride cyanate complexes was observed with various (hetero)allenes,260 alkynes,261 cyanides,261 and alkyl triflates,262 in which new NdC bonds were formed.

Me2Si

SiMe2

N

Hf

Hf

N

1 equiv CO O

Hf

Me2Si

C

1–4 atm CO

N

N H

O Hf

SiMe2

Hf

Me2Si

N

C C

N SiMe2

Hf O

or isomer Fig. 50 CO addition to N2 in a dihafnium N2 complex.

A new route toward NdC bond formation takes advantage of the partial silylation of N2 to give N2(SiR3)2 species that are electrophilic at the proximal N atom. Peters and coworkers first demonstrated this pathway with the addition of PhSiH3 to a FeNNSi2 species, which resulted in formation of a FeN(SiH2Ph)NSi2 ligand, from formal HdSi addition across the M–N bond (see Fig. 43).239 Later, they showed the migration of a hydride from Fe to an N2-derived FeNNSi2 species to form a new NdH bond. Since this reaction started from an iron(IV) hydride complex, the migration is a formal reduction, which gives an iron(II) hydrazido product.263 Most recently, it was possible to induce migration of an iron-bound phenyl ligand in a similar fashion (Fig. 51). Interestingly, this phenyl group came from benzene CdH activation, and it was possible to perform hydrocarbon activation, partial N2 silylation and migration in one pot to give a proposed cyclic mechanism for disilylaniline formation directly from N2 and benzene.234 This exciting result demonstrates the opportunities for umpolung that turns the usually nucleophilic N2 into an electrophile through silylation of the distal N. In the long term, it suggests the feasibility of creating CdN containing organic compounds catalytically from abundant hydrocarbons and N2.

Na(15c5)+ –

N

N Fe N

N

SiMe3

Na 15c5 2 Me3SiBr

N

–2 Na(15c5)Br

N

N

N

Fe

SiMe3

SiMe3 N Fe N

proposed Fig. 51 NdC bond formation to a partially silylated N2 via aryl migration.

N

N SiMe 3

Dinitrogen Binding and Functionalization

1.17.4.2.3

549

NdB bonds

Though not as common as NdSi or NdC formation, the direct formation of NdB bonds from N2 has been investigated.231 In a seminal report from Hidai, the addition of primary and secondary boranes to [trans-W(N2)(NCS)(dppe)2]− gave the corresponding boryldiazenido complexes (Fig. 52, top).264 No reactivity was observed with trans-W(N2)2(dppe)2, and the thiocyanate’s influence could come from a trans influence or from the increase in negative charge. (A)

B N Ph2 N P W P Ph2 N C S

(B)

N Ph2 N P W P Ph2 X

(nBu4)4N+ H Ph2 P P Ph2

Ar B

X

Ph2 P P Ph2

–

N H2B

X2BAr Ar = Mes, Dur, Fc X = Br, Cl R = Ph

Ph2 N P W P Ph2 N C S

Ph2 P P Ph2

N HBR2 R = Cy or R2 = C8H14

N R2 N P W P R2 N

Ph2 N P W P Ph2 N C S

N R2 P P R2

HB(C6F5)2 R = Et

Et2 N P W P Et2 H

BR2 Ph2 P P Ph2

BAr2 Et2 P P Et2

N Fig. 52 Formation of boryldiazenido complexes from various tungsten N2 species.

Increasing the electrophilicity of the borane reagent also helps to induce reactivity with N2 complexes, as trans-W(N2)2(dppe)2 undergoes 1,3-haloboration with various haloboranes to give the respective boryldiazenido complexes (Fig. 52, bottom).265 1,3-hydroboration of trans-W(N2)2(depe)2 is also possible with the use of perfluorophenylboranes, raising the possibility of combining N2 chemistry with frustrated Lewis pair chemistry.192,266 As noted above, end-on/side-on bound N2 is particularly reactive toward incoming electrophiles, and has enabled NdN cleavage using boron reagents. Two examples are shown in Fig. 53. The side-on/end-on ditantalum N2 complex (N2P)Ta (m-Z1:Z2-N2)(m-H)2Ta(N2P) (N2P ¼ PhP(CH2SiMe2NPh)2), reacts with mono- and dialkylboranes to give boryldiazenido complexes.267,268 Heating a solution of the boryldiazenido species induces cleavage of the NdN bond to yield an imido/nitride complex with concomitant loss of H2 and benzene. Interestingly, 15N labelling studies indicated that the resulting nitride comes from the NPN-pincer ligand; this was proposed to occur by silyl migration after initial loss of H2 to give a phenylimido complex, which then spontaneously loses benzene to give the final product. In a different strategy, Semproni and Chirik reported borylation of the side-on bridged dizirconium N2 complex using HBPin to give the boryldiazenido complex [(1,3,4-Me3Cp)2Zr](m-Z1:Z2-NNBPin) (m-H)[Zr(1,3,4-Me3Cp)2], and subsequent reaction with CO led to NdN bond scission and NdC bond formation to give a bridging borylimido/formamide complex.269

Fig. 53 NdN bond cleavage following borylation of N2.

550

Dinitrogen Binding and Functionalization

1.17.4.2.4

Other N–X bonds

Due to the ease of stoichiometric and catalytic silylation of N2 using silyl electrophiles, analogous reactivity of N2 complexes with germyl electrophiles was explored by Hidai and coworkers in the early 1990’s.270,271 Reactions of cis-[W(N2)2(PMe2Ph)4] with R3GeCl (R ¼ Me, Ph) in the presence of NaI led to the formation of the germyldiazenido complexes trans-[WI(NNGeR3) (PMe2Ph)4]. It was noted that these germyldiazenido complexes were much less stable than their corresponding silyl analogues, trending with the lower stability of free germyldiazenes R3GeNNR’ versus the corresponding silyldiazenes. Dinitrogen-derived NdP bond formation was inferred by Chatt and coworkers in 1970,272 but was crystallographically verified for the first time by Murahashi et al. in 2005.273 Addition of PPh2Cl to the anionic molybdenum N2 complex [Na(THF)x][((tBu) NAr)3MoN2] (Ar ¼ 3,5-Me2C6H3) led to the isolation of the phosphinodiazenido complex ((tBu)NAr)3MoNNPPh2 (Fig. 54). Further reactivity of this complex with MeOTf gave a rare diazophosphorane species, however attempts to isolate the free diazophosphorane ligand from this complex were unsuccessful. Since then, only one other example of NdP bond formation to N2 has been crystallographically verified, which is proposed to occur via N atom insertion into an FedP bond of an iron bis(phosphino)borane complex.239

Na+ N

OTf – –

N

tBu

N Ar N

tBu

Ar

Mo

PPh2 PPh2Cl

N

tBu

N

tBu

N Ar N

tBu

Ar

Ar

Mo

N N

PPh2Me MeOTf

tBu

Ar

N

tBu

N Ar N

tBu

Ar

Mo

N N

tBu

Ar

Ar = 3,5-dimethylphenyl Fig. 54 Use of a phosphorus electrophile for NdP bond formation.

1.17.5

Summary and perspectives

Several themes emerge from this overview of the binding and reactivity of N2, which are also emphasized in a recent special issue of Chemical Reviews.274 One is that N2 has a diverse set of binding modes,20 which have substantial influences on the NdN weakening, polarization, and reactivity of the complexes. Though it is tempting to imagine that there is a correlation between the ground-state weakening/reduction of the NdN bond and the ability to functionalize the N2, this is not the case: the most evident example is that the binding mode with the least weakening (terminal end-on) is the source of most stoichiometric and catalytic reactions. This is a good demonstration of the difference between ground-state stabilization (weakening) and transition-state stabilization (reactivity) that is important throughout organometallic chemistry and small-molecule activation. It is also cause for optimism, because it suggests that some N2 complexes can become reactive without requiring a massive influx of electron density into the N2 unit, and that it will eventually be possible to achieve catalytic reactions without the very strong reducing agents that predominate in the field now. Making the conversions of N2 to reduced products more energetically reasonable is a more general theme that will continue into the future. In this area, chemists are starting to focus on the ability of hydride species to provide low-energy electrons together with a driving force from release of H2.160,193 This may be part of the strategy used by natural systems to reduce N2 using nitrogenase enzymes.275 Another set of emerging concepts revolves around proton-coupled electron transfer (PCET), which also avoids the buildup of charge during the repeated steps of NdH bond formation.203 One can envision other uses for concerted mechanisms as well, for example coupling electron transfer to other positively-charged ion movement, or coupling protons/cations to the formation of other N–X bonds from N2. Hopefully, these and other strategies will enable chemists to utilize N2, the abundant resource for chemistry that is all around us.

Acknowledgments This work was supported by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences, Catalysis Program, under Award DE-SC0020315 (P.L.H.), and by a graduate fellowship from the National Science Foundation (J.E.W.), grant number DGE1752134.We thank colleagues in the Holland Group for editorial advice and assistance, particularly Alexandra Nagelski, Linda Zuckerman, Daniel Wilson, Majed Fataftah, and Erik Phipps. We thank Jonas Peters for helpful comments.

References 1. Appl, M. Ammonia. In Ullmann’s Encyclopedia of Industrial Chemistry, Wiley: Weinheim, Germany, 2011. 2. Peplow, M. Chem. World 2013, 48–51. 3. Zhang, X.; Ward, B. B.; Sigman, D. M. Chem. Rev. 2020, 120, 5308–5351.

Dinitrogen Binding and Functionalization

551

4. Lindley, B. M.; Appel, A. M.; Krogh-Jespersen, K.; Mayer, J. M.; Miller, A. J. M. ACS Energy Lett. 2016, 1, 698–704. 5. Wiig, J. A.; Hu, Y.; Lee, C. C.; Ribbe, M. W. Science 2012, 337, 1672–1675. 6. Chen, J. G.; Crooks, R. M.; Seefeldt, L. C.; Bren, K. L.; Bullock, R. M.; Darensbourg, M. Y.; Holland, P. L.; Hoffman, B.; Janik, M. J.; Jones, A. K.; Kanatzidis, M. G.; King, P.; Lancaster, K. M.; Lymar, S. V.; Pfromm, P.; Schneider, W. F.; Schrock, R. R. Science 2018, 360. eaar6611. 7. Jia, H.-P.; Quadrelli, E. A. Chem. Soc. Rev. 2014, 43, 547–564. 8. Tanabe, Y.; Nishibayashi, Y. Coord. Chem. Rev. 2013, 257, 2551–2564. 9. Holland, P. L. Dalton Trans. 2010, 39, 5415–5425. 10. Hazari, N. Chem. Soc. Rev. 2010, 39, 4044–4056. 11. MacKay, B. A.; Fryzuk, M. D. Chem. Rev. 2004, 104, 385–401. 12. Leigh, G. J. J. Organomet. Chem. 2004, 689, 3999–4005. 13. Fryzuk, M. D.; Johnson, S. A. Coord. Chem. Rev. 2000, 200-202, 379–409. 14. Hidai, M.; Mizobe, Y. Chem. Rev. 1995, 95, 1115–1133. 15. Shilov, A. E. Russ. Chem. Bull. 2003, 52, 2555–2562. 16. Minteer, S. D.; Christopher, P.; Linic, S. ACS Energy Lett. 2019, 4, 163–166. 17. Bazhenova, T. A.; Shilov, A. E. Coord. Chem. Rev. 1995, 144, 69–145. 18. Masero, F.; Perrin, M. A.; Dey, S.; Mougel, V. Chem. Eur. J. 2021, 27, 3892–3928. 19. Stucke, N.; Flöser, B. M.; Weyrich, T.; Tuczek, F. Eur. J. Inorg. Chem. 2018, 2018, 1337–1355. 20. Burford, R. J.; Fryzuk, M. D. Nat. Rev. Chem. 2017, 1, 0026. 21. Roux, Y.; Duboc, C.; Gennari, M. ChemPhysChem 2017, 18, 2606–2617. 22. Khoenkhoen, N.; de Bruin, B.; Reek, J. N. H.; Dzik, W. I. Eur. J. Inorg. Chem. 2015, 2015, 567–598. 23. Chatt, J.; Dilworth, J. R.; Richards, R. L. Chem. Rev. 1978, 78, 589–625. 24. Dilworth, J. R. Coord. Chem. Rev. 1996, 154, 163–177. 25. Shaver, M. P.; Fryzuk, M. D. Adv. Synth. Catal. 2003, 345, 1061–1076. 26. Gambarotta, S.; Scott, J. Angew. Chem. Int. Ed. 2004, 43, 5298–5308. 27. Studt, F.; Tuczek, F. J. Comput. Chem. 2006, 27, 1278–1291. 28. Hinrichsen, S.; Broda, H.; Gradert, C.; Söncksen, L.; Tuczek, F. Annu. Rep. Prog. Chem., Sect. A: Inorg. Chem. 2012, 108, 17–47. 29. Nishibayashi, Y. Inorg. Chem. 2015, 54, 9234–9247. 30. Arashiba, K.; Eizawa, A.; Tanaka, H.; Nakajima, K.; Yoshizawa, K.; Nishibayashi, Y. Bull. Chem. Soc. Japan 2017, 90, 1111–1118. 31. Liu, H.; Wei, L.; Liu, F.; Pei, Z.; Shi, J.; Wang, Z.-J.; He, D.; Chen, Y. ACS Catalysis 2019, 9, 5245–5267. 32. Chalkley, M. J.; Drover, M. W.; Peters, J. C. Chem. Rev. 2020, 120, 5582–5636. 33. Singh, D.; Buratto, W. R.; Torres, J. F.; Murray, L. J. Chem. Rev. 2020, 120, 5517–5581. 34. Coric, I.; Mercado, B. Q.; Bill, E.; Vinyard, D. J.; Holland, P. L. Nature 2015, 526, 96–99. 35. Battino, R.; Rettich, T. R.; Tominaga, T. J. Phys. Chem. Ref. Data. 1984, 13, 563–600. 36. Jabłoniec, A.; Horstmann, S.; Gmehling, J. Ind. Eng. Chem. Res. 2007, 46, 4654–4659. 37. James, R. Acta Cryst. 1948, 1, 132–134. 38. McWilliams, S. F.; Rodgers, K. R.; Lukat-Rodgers, G.; Mercado, B. Q.; Grubel, K.; Holland, P. L. Inorg. Chem. 2016, 55, 2960–2968. 39. Nakamoto, K. Infrared and Raman Spectra of Inorganic and Coordination Compounds: Theory and Applications in Inorganic Chemistry (Volume a); Wiley: New York, 1997. 40. Duan, G. Y.; Ren, Y.; Tang, Y.; Sun, Y. Z.; Chen, Y. M.; Wan, P. Y.; Yang, X. J. ChemSusChem 2020, 13, 88–96.  c, V.; Yang, S.; Bent, S. F.; 41. Nielander, A. C.; McEnaney, J. M.; Schwalbe, J. A.; Baker, J. G.; Blair, S. J.; Wang, L.; Pelton, J. G.; Andersen, S. Z.; Enemark-Rasmussen, K.; Coli Cargnello, M.; Kibsgaard, J.; Vesborg, P. C. K.; Chorkendorff, I.; Jaramillo, T. F. ACS Catal. 2019, 9, 5797–5802. 42. Hodgetts, R. Y.; Kiryutin, A. S.; Nichols, P.; Du, H.-L.; Bakker, J. M.; Macfarlane, D. R.; Simonov, A. N. ACS Energy Lett. 2020, 5, 736–741. 43. Weatherburn, M. W. Anal. Chem. 1967, 39, 971–974. 44. Chaney, A. L.; Marbach, E. P. Clin. Chem. 1962, 8, 130–132. 45. Anderson, J. S.; Rittle, J.; Peters, J. C. Nature 2013, 501, 84–87. 46. Greenlee, L. F.; Renner, J. N.; Foster, S. L. ACS Catal. 2018, 8, 7820–7827.  c, V.; Yang, S.; Schwalbe, J. A.; Nielander, A. C.; McEnaney, J. M.; Enemark-Rasmussen, K.; Baker, J. G.; Singh, A. R.; Rohr, B. A.; Statt, M. J.; Blair, S. J.; 47. Andersen, S. Z.; Coli Mezzavilla, S.; Kibsgaard, J.; Vesborg, P. C. K.; Cargnello, M.; Bent, S. F.; Jaramillo, T. F.; Stephens, I. E. L.; Nørskov, J. K.; Chorkendorff, I. Nature 2019, 570, 504–508. 48. Dabundo, R.; Lehmann, M. F.; Treibergs, L.; Tobias, C. R.; Altabet, M. A.; Moisander, P. H.; Granger, J. PLOS One 2014, 9, e110335. 49. Ohyama, T.; Kumazawa, K. Soil Sci. Plant Nutr. 1981, 27, 263–265. 50. Babcock, S. H., Jr.; Lankelma, H. P.; Vopicka, E. Inorg. Synth. 1939, 10–11. 51. Castillo, M.; Metta-Magaña, A. J.; Fortier, S. New J. Chem. 2016, 40, 1923–1926. 52. Smith, J. M.; Sadique, A. R.; Cundari, T. R.; Rodgers, K. R.; Lukat-Rodgers, G.; Lachicotte, R. J.; Flaschenriem, C. J.; Vela, J.; Holland, P. L. J. Am. Chem. Soc. 2006, 128, 756–769. 53. Millard, M. D.; Moore, C. E.; Rheingold, A. L.; Figueroa, J. S. J. Am. Chem. Soc. 2010, 132, 8921–8923. 54. Margulieux, G. W.; Semproni, S. P.; Chirik, P. J. Angew. Chem. Int. Ed. 2014, 53, 9189–9192. 55. Taube, H. Pure Appl. Chem. 1979, 51, 901–912. 56. Lindley, B. M.; Jacobs, B. P.; MacMillan, S. N.; Wolczanski, P. T. Chem. Commun. 2016, 52, 3891–3894. 57. Fryzuk, M. D. Chem. Commun. 2013, 49, 4866–4868. 58. Vaska, L.; DiLuzio, J. W. J. Am. Chem. Soc. 1961, 83, 2784–2785. 59. Collman, J. P.; Kang, J. W. J. Am. Chem. Soc. 1966, 88, 3459–3460. 60. Collman, J. P.; Kubota, M.; Sun, J.-Y.; Vastine, F. J. Am. Chem. Soc. 1967, 89, 169–170. 61. Collman, J. P.; Kubota, M.; Vastine, F. D.; Sun, J. Y.; Kang, J. W. J. Am. Chem. Soc. 1968, 90, 5430–5437. 62. Chatt, J.; Dilworth, J. R.; Leigh, G. J. J. Chem. Soc. D. 1969, 687–688. 63. Chatt, J.; Dilworth, J. R.; Leigh, G. J.; Gupta, V. D. J. Chem. Soc. A. 1971, 2631–2639. 64. Chatt, J.; Dilworth, J. R.; Leigh, G. J. J. Chem. Soc., Dalton Trans. 1973, 612–618. 65. Buhr, J. D.; Taube, H. Inorg. Chem. 1979, 18, 2208–2212. 66. Ware, D. C.; Taube, H. Inorg. Chem. 1991, 30, 4605–4610. 67. Che, C.-M.; Lam, H.-W.; Tong, W.-F.; Lai, T.-F.; Lau, T.-C. J. Chem. Sci., Chem. Commun. 1989, 1883–1884. 68. Demadis, K. D.; Meyer, T. J.; White, P. S. Inorg. Chem. 1997, 36, 5678–5679. 69. Betley, T. A.; Peters, J. C. J. Am. Chem. Soc. 2004, 126, 6252–6254. 70. Dunn, P. L.; Cook, B. J.; Johnson, S. I.; Appel, A. M.; Bullock, R. M. J. Am. Chem. Soc. 2020, 142, 17845–17858. 71. Abbenseth, J.; Finger, M.; Würtele, C.; Kasanmascheff, M.; Schneider, S. Inorg. Chem. Front. 2016, 3, 469–477. 72. Scheibel, M. G.; Askevold, B.; Heinemann, F. W.; Reijerse, E. J.; de Bruin, B.; Schneider, S. Nature Chem 2012, 4, 552–558.

552

Dinitrogen Binding and Functionalization

73. Crabtree, R. H. The Organometallic Chemistry of the Transition Metals; Wiley: New York, 2014. 74. Hunter, E. P.; Lias, S. G. In NIST Chemistry WebBook, NIST Standard Reference Database Number 69 National Institute of Standards and Technology; Linstrom, PJ, Mallard, WG, Eds.; Gaithersburg, MD, 2021, https://doi.org/10.18434/T4D303. (Retrieved August 31, 2021). 75. Dapprich, S.; Pidun, U.; Ehlers, A. W.; Frenking, G. Chem. Phys. Lett. 1995, 242, 521–526. 76. de Llarduya, J. M. M.; Villafane, F. J. Chem. Educ. 1994, 71, 480. 77. Spaeth, A. D.; Gagnon, N. L.; Dhar, D.; Yee, G. M.; Tolman, W. B. J. Am. Chem. Soc. 2017, 139, 4477–4485. 78. Klopsch, I.; Yuzik-Klimova, E. Y.; Schneider, S. Top. Organomet. Chem. 2017, 60, 71–112. 79. Burford, R. J.; Yeo, A.; Fryzuk, M. D. Coord. Chem. Rev. 2017, 334, 84–99. 80. Ohki, Y.; Fryzuk, M. D. Angew. Chem. Int. Ed. 2007, 46, 3180–3183. 81. Chirik, P. J. Dalton Trans. 2007, 16–25. 82. Chirik, P. J. Organometallics 2010, 29, 1500–1517. 83. Tanabe, Y.; Nishibayashi, Y. Coord. Chem. Rev. 2019, 381, 135–150. 84. Gardiner, M. G.; Stringer, D. N. Materials 2010, 3, 841–862. 85. Kendall, A. J.; Mock, M. T. Eur. J. Inorg. Chem. 2020, 2020, 1358–1375. 86. Murray, L. J.; Weare, W. W.; Shearer, J.; Mitchell, A. D.; Abboud, K. A. J. Am. Chem. Soc. 2014, 136, 13502–13505. 87. Zhang, S.; Fallah, H.; Gardner, E. J.; Kundu, S.; Bertke, J. A.; Cundari, T. R.; Warren, T. H. Angew. Chem. Int. Ed. 2016, 55, 9927–9931. 88. Légaré, M.-A.; Bélanger-Chabot, G.; Dewhurst, R. D.; Welz, E.; Krummenacher, I.; Engels, B.; Braunschweig, H. Science 2018, 359, 896. 89. Rösch, B.; Gentner, T. X.; Langer, J.; Färber, C.; Eyselein, J.; Zhao, L.; Ding, C.; Frenking, G.; Harder, S. Science 2021, 371, 1125–1128. 90. Peigne, B.; Aullon, G. Acta Cryst. 2015, 71, 369–386. 91. Groom, C. R.; Bruno, I. J.; Lightfoot, M. P.; Ward, S. C. Acta Cryst. B 2016, 72, 171–179. 92. Chatt, J.; Crabtree, R. H.; Richards, R. L. J. Chem. Soc., Dalton Trans. 1972, 534. 93. Chatt, J.; Crabtree, R. H.; Jeffery, E. A.; Richards, R. L. J. Chem. Soc., Dalton Trans. 1973, 1167–1172. 94. Connor, G. P.; Holland, P. L. Catal. Today 2017, 286, 21–40. 95. Shanahan, J. P.; Szymczak, N. K. J. Am. Chem. Soc. 2019, 141, 8550–8556. 96. Geri, J. B.; Shanahan, J. P.; Szymczak, N. K. J. Am. Chem. Soc. 2017, 139, 5952–5956. 97. Ferguson, R.; Solari, E.; Floriani, C.; Osella, D.; Ravera, M.; Re, N.; Chiesi-Villa, A.; Rizzoli, C. J. Am. Chem. Soc. 1997, 119, 10104–10115. 98. Goettker-Schnetmann, I.; White, P. S.; Brookhart, M. Organometallics 2004, 23, 1766–1776. 99. Mercer, M.; Crabtree, R. H.; Richards, R. L. J. Chem. Soc., Chem. Commun. 1973, 808–809. 100. Ding, K.; Pierpont, A. W.; Brennessel, W. W.; Lukat-Rodgers, G.; Rodgers, K. R.; Cundari, T. R.; Bill, E.; Holland, P. L. J. Am. Chem. Soc. 2009, 131, 9471–9472. 101. Pierpont, A. W.; Cundari, T. R. J. Coord. Chem. 2011, 64, 3123–3135. 102. Yamout, L. S.; Ataya, M.; Hasanayn, F.; Holland, P. L.; Miller, A. J. M.; Goldman, A. S. J. Am. Chem. Soc. 2021, 143, 9744–9757. 103. Ertl, G.; Lee, S. B.; Weiss, M. Surface Sci. 1982, 114, 515–526. 104. Grunze, M.; Strasser, G.; Golze, M. Appl. Phys. A. 1987, 44, 19–29. 105. Mortensen, J. J.; Hansen, L. B.; Hammer, B.; Nørskov, J. K. J. Catal. 1999, 182, 479–488. 106. Christoffersen, E.; Mortensen, J.-J.; Stoltze, P.; Nørskov, J. K. Isr. J. Chem. 1998, 38, 279–284. 107. Grubel, K.; Brennessel, W. W.; Mercado, B. Q.; Holland, P. L. J. Am. Chem. Soc. 2014, 136, 16807–16816. 108. Sorsche, D.; Miehlich, M. E.; Searles, K.; Gouget, G.; Zolnhofer, E. M.; Fortier, S.; Chen, C.-H.; Gau, M.; Carroll, P. J.; Murray, C. B.; Caulton, K. G.; Khusniyarov, M. M.; Meyer, K.; Mindiola, D. J. J. Am. Chem. Soc. 2020, 142, 8147–8159. 109. Liu, T.; Gau, M. R.; Tomson, N. C. J. Am. Chem. Soc. 2020, 142, 8142–8146. 110. Cusanelli, A.; Sutton, D. Organometallics 1996, 15, 1457–1464. 111. Fomitchev, D. V.; Bagley, K. A.; Coppens, P. J. Am. Chem. Soc. 2000, 122, 532–533. 112. Schaniel, D.; Woike, T.; Delley, B.; Boskovic, C.; Güdel, H.-U. Phys. Chem. Chem. Phys. 2008, 10, 5531–5538. 113. Morello, L.; Yu, P.; Carmichael, C. D.; Patrick, B. O.; Fryzuk, M. D. J. Am. Chem. Soc. 2005, 127, 12796–12797. 114. Schmiege, B. M.; Ziller, J. W.; Evans, W. J. Inorg. Chem. 2010, 49, 10506–10511. 115. Lv, Z.-J.; Huang, Z.; Zhang, W.-X.; Xi, Z. J. Am. Chem. Soc. 2019, 141, 8773–8777. 116. Evans, W. J.; Lee, D. S.; Rego, D. B.; Perotti, J. M.; Kozimor, S. A.; Moore, E. K.; Ziller, J. W. J. Am. Chem. Soc. 2004, 126, 14574–14582. 117. Evans, W. J.; Lee, D. S.; Ziller, J. W. J. Am. Chem. Soc. 2004, 126, 454–455. 118. Evans, W. J.; Fang, M.; Zucchi, G.; Furche, F.; Ziller, J. W.; Hoekstra, R. M.; Zink, J. I. J. Am. Chem. Soc. 2009, 131, 11195–11202. 119. Elwell, C. E.; Gagnon, N. L.; Neisen, B. D.; Dhar, D.; Spaeth, A. D.; Yee, G. M.; Tolman, W. B. Chem. Rev. 2017, 117, 2059–2107. 120. Pool, J. A.; Bernskoetter, W. H.; Chirik, P. J. J. Am. Chem. Soc. 2004, 126, 14326–14327. 121. Pool, J. A.; Lobkovsky, E.; Chirik, P. J. Nature 2004, 427, 527–530. 122. de Wolf, J. M.; Blaauw, R.; Meetsma, A.; Teuben, J. H.; Gyepes, R.; Varga, V.; Mach, K.; Veldman, N.; Spek, A. L. Organometallics 1996, 15, 4977–4983. 123. Hanna, T. E.; Bernskoetter, W. H.; Bouwkamp, M. W.; Lobkovsky, E.; Chirik, P. J. Organometallics 2007, 26, 2431–2438. 124. Semproni, S. P.; Milsmann, C.; Chirik, P. J. Organometallics 2012, 31, 3672–3682. 125. Manriquez, J. M.; Bercaw, J. E. J. Am. Chem. Soc. 1974, 96, 6229–6230. 126. Roddick, D. M.; Fryzuk, M. D.; Seidler, P. F.; Hillhouse, G. L.; Bercaw, J. E. Organometallics 1985, 4, 97–104. 127. Semproni, S. P.; Knobloch, D. J.; Milsmann, C.; Chirik, P. J. Angew. Chem. Int. Ed. 2013, 52, 5372–5376. 128. Fryzuk, M. D.; Johnson, S. A.; Rettig, S. J. J. Am. Chem. Soc. 1998, 120, 11024–11025. 129. Fryzuk, M. D.; Johnson, S. A.; Patrick, B. O.; Albinati, A.; Mason, S. A.; Koetzle, T. F. J. Am. Chem. Soc. 2001, 123, 3960–3973. 130. Pun, D.; Lobkovsky, E.; Chirik, P. J. J. Am. Chem. Soc. 2008, 130, 6047–6054. 131. Yeo, A.; Shaver, M. P.; Fryzuk, M. D. Z. Anorg. Allg. Chem. 2015, 641, 123–127. 132. Mo, Z.; Shima, T.; Hou, Z. Angew. Chem. Int. Ed. 2020, 59, 8635–8644. 133. Wang, B.; Luo, G.; Nishiura, M.; Hu, S.; Shima, T.; Luo, Y.; Hou, Z. J. Am. Chem. Soc. 2017, 139, 1818–1821. 134. Pez, G. P.; Apgar, P.; Crissey, R. K. J. Am. Chem. Soc. 2002, 104, 482–490. 135. Eaton, M. C.; Catalano, V. J.; Shearer, J.; Murray, L. J. J. Am. Chem. Soc. 2021, 143, 5649–5653. 136. Guillemot, G.; Castellano, B.; Prangé, T.; Solari, E.; Floriani, C. Inorg. Chem. 2007, 46, 5152–5154. 137. Dube, T.; Conochi, S.; Gambarotta, S.; Yap, G. P. A.; Vasapollo, G. Angew. Chem. Int. Ed. 1999, 38, 3657–3659. 138. Dube, T.; Ganesan, M.; Conoci, S.; Gambarotta, S.; Yap, G. P. A. Organometallics 2000, 19, 3716–3721. 139. Bérubé, C. D.; Yazdanbakhsh, M.; Gambarotta, S.; Yap, G. P. A. Organometallics 2003, 22, 3742–3747. 140. Shan, H.; Yang, Y.; James, A. J.; Sharp, P. R. Science 1997, 275, 1460–1462. 141. Chisholm, M. H.; Cotton, F. A.; Frenz, B. A.; Reichert, W. W.; Shive, L. W.; Stults, B. R. J. Am. Chem. Soc. 1976, 98, 4469–4476. 142. Laplaza, C. E.; Cummins, C. C. Science 1995, 268, 861–863. 143. Laplaza, C. E.; Johnson, M. J. A.; Peters, J.; Odom, A. L.; Kim, E.; Cummins, C. C.; George, G. N.; Pickering, I. J. J. Am. Chem. Soc. 1996, 118, 8623–8638.

Dinitrogen Binding and Functionalization

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. 191. 192. 193. 194. 195. 196. 197. 198. 199. 200. 201. 202. 203. 204. 205. 206. 207. 208. 209. 210. 211. 212. 213. 214.

553

Peters, J. C.; Cherry, J.-P. F.; Thomas, J. C.; Baraldo, L.; Mindiola, D. J.; Davis, W. M.; Cummins, C. C. J. Am. Chem. Soc. 1999, 121, 10053–10067. Hanna, T. E.; Lobkovsky, E.; Chirik, P. J. Organometallics 2009, 28, 4079–4088. Pool, J. A.; Chirik, P. J. Can. J. Chem. 2005, 83, 286–295. McWilliams, S. F.; Holland, P. L. Acc. Chem. Res. 2015, 48, 2059–2065. Smith, J. M.; Lachicotte, R. J.; Pittard, K. A.; Cundari, T. R.; Lukat-Rodgers, G.; Rodgers, K. R.; Holland, P. L. J. Am. Chem. Soc. 2001, 123, 9222–9223. Bellows, S. M.; Arnet, N. A.; Gurubasavaraj, P. M.; Brennessel, W. W.; Bill, E.; Cundari, T. R.; Holland, P. L. J. Am. Chem. Soc. 2016, 138, 12112–12123. Rodriguez, M. M.; Bill, E.; Brennessel, W. W.; Holland, P. L. Science 2011, 334, 780–783. Figg, T. M.; Holland, P. L.; Cundari, T. R. Inorg. Chem. 2012, 51, 7546–7550. Sato, M.; Tatsumi, T.; Kodama, T.; Hidai, M.; Uchida, T.; Uchida, Y. J. Am. Chem. Soc. 1978, 100, 4447–4452. Luo, X.-L.; Kubas, G. J.; Burns, C. J.; Eckert, J. Inorg. Chem. 1994, 33, 5219–5229. Kirillov, A. M.; Haukka, M.; Guedes da Silva, M. F. C.; Fraústo da Silva, J. J. R.; Pombeiro, A. J. L. J. Organomet. Chem. 2006, 691, 4153–4158. Lee, J. P.; Ke, Z.; Ramírez, M. A.; Gunnoe, T. B.; Cundari, T. R.; Boyle, P. D.; Petersen, J. L. Organometallics 2009, 28, 1758–1775. Drance, M. J.; Sears, J. D.; Mrse, A. M.; Moore, C. E.; Rheingold, A. L.; Neidig, M. L.; Figueroa, J. S. Science 2019, 363, 1203–1205. Kandler, H.; Gauss, C.; Bidell, W.; Rosenberger, S.; Buergi, T.; Eremenko, I. L.; Veghini, D.; Orama, O.; Burger, P.; Berke, H. Chem. Eur. J. 1995, 1, 541–548. Bruch, Q. J.; Lindley, B. M.; Askevold, B.; Schneider, S.; Miller, A. J. M. Inorg. Chem. 2018, 57, 1964–1975. Creutz, S. E.; Peters, J. C. J. Am. Chem. Soc. 2014, 136, 1105–1115. Del Castillo, T. J.; Thompson, N. B.; Peters, J. C. J. Am. Chem. Soc. 2016, 138, 5341–5350. Matson, B. D.; Peters, J. C. ACS Catal. 2018, 8, 1448–1455. Fajardo, J.; Peters, J. C. Inorg. Chem. 2021, 60, 1220–1227. Forrest, S. J. K.; Schluschaß, B.; Yuzik-Klimova, E. Y.; Schneider, S. Chem. Rev. 2021, 121, 6522–6587. Cui, Q.; Musaev, D. G.; Svensson, M.; Sieber, S.; Morokuma, K. J. Am. Chem. Soc. 1995, 117, 12366–12367. Bruch, Q. J.; Connor, G. P.; McMillion, N. D.; Goldman, A. S.; Hasanayn, F.; Holland, P. L.; Miller, A. J. M. ACS Catal. 2020, 10, 10826–10846. Solari, E.; Da Silva, C.; Iacono, B.; Hesschenbrouck, J.; Rizzoli, C.; Scopelliti, R.; Floriani, C. Angew. Chem. Int. Ed. 2001, 40, 3907–3909. Curley, J. J.; Cook, T. R.; Reece, S. Y.; Muller, P.; Cummins, C. C. J. Am. Chem. Soc. 2008, 130, 9394–9405. Bruch, Q. J.; Connor, G. P.; Chen, C. H.; Holland, P. L.; Mayer, J. M.; Hasanayn, F.; Miller, A. J. M. J. Am. Chem. Soc. 2019, 141, 20198–20208. Miyazaki, T.; Tanaka, H.; Tanabe, Y.; Yuki, M.; Nakajima, K.; Yoshizawa, K.; Nishibayashi, Y. Angew. Chem. Int. Ed. 2014, 53, 11488–11492. Falcone, M.; Chatelain, L.; Scopelliti, R.; Živkovic, I.; Mazzanti, M. Nature 2017, 547, 332–335. Keane, A. J.; Yonke, B. L.; Hirotsu, M.; Zavalij, P. Y.; Sita, L. R. J. Am. Chem. Soc. 2014, 136, 9906–9909. Ishida, Y.; Kawaguchi, H. J. Am. Chem. Soc. 2014, 136, 16990–16993. Shima, T.; Yang, J.; Luo, G.; Luo, Y.; Hou, Z. J. Am. Chem. Soc. 2020, 142, 9007–9016. Sivasankar, C.; Madarasi, P. K.; Tamizmani, M. Eur. J. Inorg. Chem. 2020, 2020, 1383–1395. Chatt, J.; Heath, G. A.; Richards, R. L. J. Chem. Soc., Chem. Commun. 1972, 1010–1011. Chatt, J.; Heath, G. A.; Richards, R. L. J. Chem. Soc., Dalton Trans. 1974, 2074–2082. Chatt, J.; Pearman, A. J.; Richards, R. L. Nature 1975, 253, 39–40. Chatt, J.; Pearman, A. J.; Richards, R. L. J. Chem. Soc., Dalton Trans. 1976, 1520–1524. Chatt, J.; Pearman, A. J.; Richards, R. L. J. Chem. Soc., Dalton Trans. 1977, 1852–1860. Takahashi, T.; Mizobe, Y.; Sato, M.; Uchida, Y.; Hidai, M. J. Am. Chem. Soc. 1980, 102, 7461–7467. Baumann, J. A.; George, T. A. J. Am. Chem. Soc. 1980, 102, 6153–6154. George, T. A.; Tisdale, R. C. J. Am. Chem. Soc. 1985, 107, 5157–5159. Leigh, G. J.; Jimenez-Tenorio, M. J. Am. Chem. Soc. 2002, 113, 5862–5863. Hills, A.; Hughes, D. L.; Jimenez-Tenorio, M.; Leigh, G. J.; Rowley, A. T. J. Chem. Soc., Dalton Trans. 1993, 3041. Hall, D. A.; Leigh, G. J. J. Chem. Soc., Dalton Trans. 1996, 3539. Gilbertson, J. D.; Szymczak, N. K.; Tyler, D. R. J. Am. Chem. Soc. 2005, 127, 10184–10185. Doyle, L. R.; Hill, P. J.; Wildgoose, G. G.; Ashley, A. E. Dalton Trans. 2016, 45, 7550–7554. Manriquez, J. M.; McAlister, D. R.; Rosenberg, E.; Shiller, A. M.; Williamson, K. L.; Chan, S. I.; Bercaw, J. E. J. Am. Chem. Soc. 1978, 100, 3078–3083. Manriquez, J. M.; Sanner, R. D.; Marsh, R. E.; Bercaw, J. E. J. Am. Chem. Soc. 1976, 98, 3042–3044. Tanabe, Y.; Nishibayashi, Y. Chem. Soc. Rev. 2021, 50, 5201–5242. McWilliams, S. F.; Bill, E.; Lukat-Rodgers, G.; Rodgers, K. R.; Mercado, B. Q.; Holland, P. L. J. Am. Chem. Soc. 2018, 140, 8586–8598. Simonneau, A.; Turrel, R.; Vendier, L.; Etienne, M. Angew. Chem. Int. Ed. 2017, 56, 12268–12272. Ballmann, J.; Munha, R. F.; Fryzuk, M. D. Chem. Commun. 2010, 46, 1013–1025. Jia, G.; Morris, R. H.; Schweitzer, C. T. Inorg. Chem. 1991, 30, 593–594. Nishibayashi, Y.; Iwai, S.; Hidai, M. Science 1998, 279, 540–542. Fryzuk, M. D.; Love, J. B.; Rettig, S. J.; Young, V. G. Science 1997, 275, 1445–1447. Cohen, J. D.; Fryzuk, M. D.; Loehr, T. M.; Mylvaganam, M.; Rettig, S. J. Inorg. Chem. 1998, 37, 112–119. Hidai, M.; Takahashi, T.; Yokotake, I.; Uchida, Y. Chem. Lett. 1980, 9, 645–646. Nishihara, H.; Mori, T.; Tsurita, Y.; Nakano, K.; Saito, T.; Sasaki, Y. J. Am. Chem. Soc. 1982, 104, 4367–4372. Berry, J. F. Comments Inorg. Chem. 2009, 30, 28–66. Smith, J. M. Prog. Inorg. Chem. 2014, 58, 417–470. Chalkley, M. J.; Peters, J. C. Eur. J. Inorg. Chem. 2020, 2020, 1353–1357. Chalkley, M. J.; Del Castillo, T. J.; Matson, B. D.; Roddy, J. P.; Peters, J. C. ACS Central Sci. 2017, 3, 217–223. Rittle, J.; Peters, J. C. J. Am. Chem. Soc. 2017, 139, 3161–3170. Bezdek, M. J.; Chirik, P. J. Dalton Trans. 2016, 45, 15922–15930. Scheibel, M. G.; Abbenseth, J.; Kinauer, M.; Heinemann, F. W.; Würtele, C.; de Bruin, B.; Schneider, S. Inorg. Chem. 2015, 54, 9290–9302. Bezdek, M. J.; Guo, S.; Chirik, P. J. Science 2016, 354, 730. Bezdek, M. J.; Chirik, P. J. Angew. Chem. Int. Ed. 2018, 57, 2224–2228. Bhattacharya, P.; Heiden, Z. M.; Wiedner, E. S.; Raugei, S.; Piro, N. A.; Kassel, W. S.; Bullock, R. M.; Mock, M. T. J. Am. Chem. Soc. 2017, 139, 2916–2919. Johnson, S. I.; Heins, S. P.; Klug, C. M.; Wiedner, E. S.; Bullock, R. M.; Raugei, S. Chem. Commun. 2019, 55, 5083–5086. Bhattacharya, P.; Heiden, Z. M.; Chambers, G. M.; Johnson, S. I.; Bullock, R. M.; Mock, M. T. Angew. Chem. Int. Ed. Engl. 2019, 58, 11618–11624. Najafian, A.; Cundari, T. R. J. Phys. Chem. A 2019, 123, 7973–7982. Dunn, P. L.; Johnson, S. I.; Kaminsky, W.; Bullock, R. M. J. Am. Chem. Soc. 2020, 142, 3361–3365. Bezdek, M. J.; Pappas, I.; Chirik, P. J. Determining and understanding N-H bond strengths in synthetic nitrogen fixation cycles. In Nitrogen Fixation; Nishibayashi, Y., Ed.; Springer International Publishing: Cham, 2017;; pp 1–21. 215. Wang, D.; Loose, F.; Chirik, P. J.; Knowles, R. R. J. Am. Chem. Soc. 2019, 141, 4795–4799.

554

216. 217. 218. 219. 220. 221. 222. 223. 224. 225. 226. 227. 228. 229. 230. 231. 232. 233. 234. 235. 236. 237. 238. 239. 240. 241. 242. 243. 244. 245. 246. 247. 248. 249. 250. 251. 252. 253. 254. 255. 256. 257. 258. 259. 260. 261. 262. 263. 264. 265. 266. 267. 268. 269. 270. 271. 272. 273. 274. 275.

Dinitrogen Binding and Functionalization

Stephan, G. C.; Sivasankar, C.; Studt, F.; Tuczek, F. Chem. Eur. J. 2008, 14, 644–652. Pappas, I.; Chirik, P. J. J. Am. Chem. Soc. 2016, 138, 13379–13389. Warren, J. J.; Tronic, T. A.; Mayer, J. M. Chem. Rev. 2010, 110, 6961–7001. Yandulov, D. V.; Schrock, R. R. Inorg. Chem. 2005, 44, 1103–1117. Nesbit, M. A.; Oyala, P. H.; Peters, J. C. J. Am. Chem. Soc. 2019, 141, 8116–8127. Kaljurand, I.; Kutt, A.; Soovali, L.; Rodima, T.; Maemets, V.; Leito, I.; Koppel, I. A. J. Org. Chem. 2005, 70, 1019–1028. Garrido, G.; Rosés, M.; Ràfols, C.; Bosch, E. J. Solution Chem. 2008, 37, 689–700. Wise, C. F.; Agarwal, R. G.; Mayer, J. M. J. Am. Chem. Soc. 2020, 142, 10681–10691. Tyburski, R.; Liu, T.; Glover, S. D.; Hammarstrom, L. J. Am. Chem. Soc. 2021, 143, 560–576. Chalkley, M. J.; Del Castillo, T. J.; Matson, B. D.; Peters, J. C. J. Am. Chem. Soc. 2018, 140, 6122–6129. Ashida, Y.; Arashiba, K.; Nakajima, K.; Nishibayashi, Y. Nature 2019, 568, 536–540. Hwu, J. R.; Wetzel, J. M. J. Org. Chem. 1985, 50, 3946–3948. Hwu, J. R.; Wang, N. Chem. Rev. 1989, 89, 1599–1615. Fleming, I. Chem. Soc. Rev. 1981, 10, 83–111. Lee, Y.; Mankad, N. P.; Peters, J. C. Nat. Chem. 2010, 2, 558–565. Kim, S.; Loose, F.; Chirik, P. J. Chem. Rev. 2020, 120, 5637–5681. Miyazaki, T.; Tanabe, Y.; Yuki, M.; Miyake, Y.; Nakajima, K.; Nishibayashi, Y. Chem. Eur. J. 2013, 19, 11874–11877. Yin, J.; Li, J.; Wang, G.-X.; Yin, Z.-B.; Zhang, W.-X.; Xi, Z. J. Am. Chem. Soc. 2019, 141, 4241–4247. McWilliams, S. F.; Broere, D. L. J.; Halliday, C. J. V.; Bhutto, S. M.; Mercado, B. Q.; Holland, P. L. Nature 2020, 584, 221–226. Oshita, H.; Mizobe, Y.; Hidai, M. Organometallics 1992, 11, 4116–4123. Moret, M.-E.; Peters, J. C. J. Am. Chem. Soc. 2011, 133, 18118–18121. Tanaka, H.; Sasada, A.; Kouno, T.; Yuki, M.; Miyake, Y.; Nakanishi, H.; Nishibayashi, Y.; Yoshizawa, K. J. Am. Chem. Soc. 2011, 133, 3498–3506. Semproni, S. P.; Lobkovsky, E.; Chirik, P. J. J. Am. Chem. Soc. 2011, 133, 10406–10409. Suess, D. L. M.; Peters, J. C. J. Am. Chem. Soc. 2013, 135, 4938–4941. Tanabe, Y.; Nishibayashi, Y. Coord. Chem. Rev. 2019, 389, 73–93. Kuriyama, S.; Arashiba, K.; Nakajima, K.; Matsuo, Y.; Tanaka, H.; Ishii, K.; Yoshizawa, K.; Nishibayashi, Y. Nature Commun. 2016, 7, 12181. Imayoshi, R.; Nakajima, K.; Takaya, J.; Iwasawa, N.; Nishibayashi, Y. Eur. J. Inorg. Chem. 2017, 2017, 3769–3778. Kuriyama, S.; Arashiba, K.; Nakajima, K.; Tanaka, H.; Yoshizawa, K.; Nishibayashi, Y. Eur. J. Inorg. Chem. 2016, 2016, 4856–4861. Kawakami, R.; Kuriyama, S.; Tanaka, H.; Arashiba, K.; Konomi, A.; Nakajima, K.; Yoshizawa, K.; Nishibayashi, Y. Chem. Commun. 2019, 55, 14886–14889. Meng, F.; Kuriyama, S.; Tanaka, H.; Egi, A.; Yoshizawa, K.; Nishibayashi, Y. Angew. Chem. Int. Ed. 2021, 60, 13906. Kawakami, R.; Kuriyama, S.; Tanaka, H.; Konomi, A.; Yoshizawa, K.; Nishibayashi, Y. Chem. Lett. 2020, 49, 794–797. Yuki, M.; Tanaka, H.; Sasaki, K.; Miyake, Y.; Yoshizawa, K.; Nishibayashi, Y. Nat. Commun. 2012, 3, 1254. Imayoshi, R.; Tanaka, H.; Matsuo, Y.; Yuki, M.; Nakajima, K.; Yoshizawa, K.; Nishibayashi, Y. Chem. Eur. J. 2015, 21, 8905–8909. Siedschlag, R. B.; Bernales, V.; Vogiatzis, K. D.; Planas, N.; Clouston, L. J.; Bill, E.; Gagliardi, L.; Lu, C. C. J. Am. Chem. Soc. 2015, 137, 4638–4641. Lv, Z.-J.; Wei, J.; Zhang, W.-X.; Chen, P.; Deng, D.; Shi, Z.-J.; Xi, Z. Nat. Sci. Rev. 2020, 7, 1564–1583. Yoshida, T.; Adachi, T.; Ueda, T.; Kaminaka, M.; Sasaki, N.; Higuchi, T.; Aoshima, T.; Mega, I.; Mizobe, Y.; Hidai, M. Angew. Chem. Int. Ed. Engl. 1989, 28, 1040–1042. Ishii, Y.; Ishino, Y.; Aoki, T.; Hidai, M. J. Am. Chem. Soc. 1992, 114, 5429–5430. Ishii, Y.; Kawaguchi, M.; Ishino, Y.; Aoki, T.; Hidai, M. Organometallics 1994, 13, 5062–5071. Sellmann, D.; Weiss, W. Angew. Chem. 1977, 89, 918–919. Sellmann, D.; Weiss, W. J. Organomet. Chem. 1978, 160, 183–196. Sellmann, D.; Weiss, W. Angew. Chem. 1978, 90, 295–296. Nakanishi, Y.; Ishida, Y.; Kawaguchi, H. Angew. Chem. Int. Ed. 2017, 56, 9193–9197. Morello, L.; Love, J. B.; Patrick, B. O.; Fryzuk, M. D. J. Am. Chem. Soc. 2004, 126, 9480–9481. Knobloch, D. J.; Lobkovsky, E.; Chirik, P. J. Nat. Chem. 2009, 2, 30–35. Semproni, S. P.; Chirik, P. J. Organometallics 2014, 33, 3727–3737. Semproni, S. P.; Chirik, P. J. J. Am. Chem. Soc. 2013, 135, 11373–11383. Semproni, S. P.; Chirik, P. J. Angew. Chem. Int. Ed. 2013, 52, 12965–12969. Deegan, M. M.; Peters, J. C. Chem. Sci. 2018, 9, 6264–6270. Ishino, H.; Ishii, Y.; Hidai, M. Chem. Lett. 1998, 677–678. Rempel, A.; Mellerup, S. K.; Fantuzzi, F.; Herzog, A.; Deißenberger, A.; Bertermann, R.; Engels, B.; Braunschweig, H. Chem. Eur. J. 2020, 26, 16019–16027. Coffinet, A.; Specklin, D.; Vendier, L.; Etienne, M.; Simonneau, A. Chem. Eur. J. 2019, 25, 14300–14303. Fryzuk, M. D.; MacKay, B. A.; Johnson, S. A.; Patrick, B. O. Angew. Chem. Int. Ed. 2002, 41, 3709–3712. MacKay, B. A.; Johnson, S. A.; Patrick, B. O.; Fryzuk, M. D. Can. J. Chem. 2005, 83, 315–323. Semproni, S. P.; Chirik, P. J. Eur. J. Inorg. Chem. 2013, 2013, 3907–3915. Oshita, H.; Mizobe, Y.; Hidai, M. Chem. Lett. 1990, 1303–1306. Oshita, H.; Mizobe, Y.; Hidai, M. J. Organomet. Chem. 1993, 456, 213–220. Chatt, J.; Dilworth, J. R.; Leigh, G. J.; Richards, R. L. J. Chem. Soc. D 1970, 955–956. Murahashi, T.; Clough, C. R.; Figueroa, J. S.; Cummins, C. C. Angew. Chem. Int. Ed. 2005, 44, 2560–2563. Holland, P. L. Chem. Rev. 2020, 120, 4919–4920. Lukoyanov, D. A.; Yang, Z.-Y.; Dean, D. R.; Seefeldt, L. C.; Raugei, S.; Hoffman, B. M. J. Am. Chem. Soc. 2020, 142, 21679–21690.

1.18

Lewis Acid Participation in Organometallic Chemistry

Julia B Curley, Nilay Hazari, and Tanya M Townsend, Department of Chemistry, Yale University, New Haven, CT, United States © 2022 Elsevier Ltd. All rights reserved.

1.18.1 1.18.2 1.18.2.1 1.18.2.2 1.18.2.3 1.18.2.4 1.18.2.5 1.18.2.6 1.18.2.7 1.18.3 1.18.3.1 1.18.3.2 1.18.3.3 1.18.3.4 1.18.4 1.18.4.1 1.18.4.1.1 1.18.4.1.2 1.18.5 1.18.5.1 1.18.5.1.1 1.18.5.1.2 1.18.5.1.3 1.18.5.2 1.18.5.2.1 1.18.5.2.2 1.18.6 1.18.6.1 1.18.6.1.1 1.18.6.1.2 1.18.6.1.3 1.18.6.2 1.18.6.2.1 1.18.6.2.2 1.18.6.3 1.18.6.3.1 1.18.7 1.18.7.1 1.18.7.2 1.18.7.3 1.18.7.4 1.18.7.5 1.18.8 1.18.8.1 1.18.8.2 1.18.8.3 1.18.9 1.18.9.1 1.18.9.2 1.18.10 References

Introduction Oxidative addition Stoichiometric oxidative addition Hydrocyanation Hydroarylation and hydroalkylation Carbocyanation Cross-coupling Cycloaddition Carbonylation Reductive elimination Stoichiometric reductive elimination Hydrocyanation Decarbonylation Cross-coupling 1,1-Insertion CO insertion Stoichiometric CO insertion CO hydrogenation 1,2-Insertion CO2 insertion Stoichiometric CO2 insertion CO2 reduction Formic acid or methanol dehydrogenation Carbonyl insertion Ester or carboxylic acid hydrogenation Amide hydrogenation to amines Ligand addition or abstraction Alkyl abstraction Stoichiometric alkyl abstraction Olefin polymerization Olefin trimerization Halogen abstraction or exchange Stoichiometric halogen abstraction from metal bound CdX bonds Arylation of amides Hydride abstraction Stoichiometric hydride abstraction for hydricity determination Ligand substitution Stoichiometric ligand substitution Olefin isomerization Carbonylation Alkene hydrogenation CO2 hydrogenation b-Hydride elimination Stoichiometric b-hydride elimination Acrylate synthesis Allene carboxylation 1,2-Addition Stoichiometric 1,2-addition H2 activation for hydrogenation Conclusions

Comprehensive Organometallic Chemistry IV

https://doi.org/10.1016/B978-0-12-820206-7.00010-X

556 557 557 558 559 559 560 560 561 562 562 564 564 564 565 565 565 566 566 566 566 567 568 569 569 570 570 571 571 571 571 572 572 573 573 573 573 574 574 575 575 576 577 577 578 578 578 578 579 580 580

555

556

Lewis Acid Participation in Organometallic Chemistry

Abbreviations Ar − BArF 4 DFT DMF LA MAD NMR OA OM OTf RE TM TOF TON triphos TS

1.18.1

Aryl Tetrakis[3,5-bis(trifluoromethyl)phenyl]borate Density functional theory (N,N-dimethylformamide) Lewis acid Methylaluminum bis(2,6-di-tert-butyl-4-methylphenoxide) Nuclear magnetic resonance Oxidative addition Organometallic Trifluoromethanesulfonate Reductive elimination Transition metal Turnover frequency Turnover number Bis(diphenylphosphinoethyl)methylphosphine Transition state

Introduction

G. N. Lewis defined acids (now called Lewis acids, LAs) as compounds with an empty orbital that can accept a pair of electrons.1 This general definition of acidity expanded the Brönsted classification, which limited acids to compounds that could transfer a proton.2 Over time, LAs have become ubiquitous as additives and co-catalysts in organometallic (OM) chemistry and often enable reactivity that is not possible by using a metal complex alone. A range of LAs are now commonly utilized to facilitate elementary OM reactions, such as oxidative addition (OA) and reductive elimination (RE), including main group compounds such as boranes, alkali metal salts such as Li+, and lanthanide salts such as those containing Yb3+ cations. Exogenous LAs can enhance both catalytic and stoichiometric OM reactions in a variety of ways (Scheme 1A). For example, LAs can be used as reagents in stoichiometric OM reactions, and as both stoichiometric and catalytic additives in OM-catalyzed reactions. Specific examples of LAs acting in these roles are provided in this article. LAs are also frequently utilized in tandem catalytic systems where an OM catalyst facilitates one (or more) steps in a reaction cascade and a LA facilitates other steps, but this is beyond the scope of this article and will not be discussed herein.

(A)

(B)

Scheme 1 (A) General classes of OM reactions involving LAs, as organized in this article. (B) Examples of LAs bound covalently to OM complexes.

In addition to being used as exogenous additives, LAs can covalently bind to metal centers, either as Z-type ligands or pendant groups on traditional donor ligands (Scheme 1B). Z-type ligands are s-acceptors with an empty orbital that accepts two electrons from the metal center upon binding.3 Evidence of LA coordination to metal centers was first reported in the 1970s and it is now relatively common to incorporate s-accepting groups into chelating ligands to stabilize the resulting metal complex.3c,4 This class of ligand imparts novel reactivity to OM complexes which is discussed throughout this article. LAs are also frequently appended to traditional X- and L-type ligands to generate ambiphilic species (Scheme 1B).5 The Lewis acidic moiety can remain pendant or interact and form an adduct with another ligand on the metal center. Tethering the LA to the metal center in this fashion ensures a high local concentration of the LA to promote elementary OM reactions, which is also described here.

Lewis Acid Participation in Organometallic Chemistry

557

Understanding the differences in affinities of LAs for electron pairs is important for the rational selection of LAs for specific applications. Unfortunately, unlike the strengths of Brönsted acids, which can be readily quantified using the pKa scale, the relative strengths of Lewis acids are more difficult to quantify, in part because of the large variations in the steric and electronic properties of LAs.6 There are, however, a number of methods that are utilized for quantifying Lewis acidity within classes of compounds. One simple strategy, adapted from the Gutmann solvent acceptor number scale, measures the change in the 31P Nuclear Magnetic Resonance (NMR) chemical shift of Et3PO upon complexation with a LA.7 This method is typically applied to boron-based LAs, whose acidity can be increased through the introduction of electron withdrawing substituents. Another strategy to measure Lewis acidity is based on the enthalpy of adduct formation between a LA and an OM carbonyl or cyano complex, which is quantified using anaerobic solution calorimetry.8 When it is not possible to use experimental methods to assess Lewis acidity, computational techniques can be utilized. One popular method is to determine the affinity of a LA for the fluoride ion.9 The high basicity and small size of the fluoride ion ensure that it binds to virtually any LA, and consequently this is a highly generalizable approach to quantifying Lewis acidity. In this article, we organize the applications of LAs in OM chemistry by the elementary reaction step in which the LA participates. For each reaction, we first describe examples of a LA facilitating stoichiometric OM reactions, before discussing the potential role of LAs in enhancing the relevant elementary reaction in catalysis. The elementary reactions discussed include OA (Section 1.18.2), RE (Section 1.18.3), 1,1-insertion (Section 1.18.4), 1,2-insertion (Section 1.18.5), ligand addition or abstraction (Section 1.18.6), ligand substitution (Section 1.18.7), b-hydride elimination (Section 1.18.8), and 1,2-addition (Section 1.18.9). Examples of cases where LAs are utilized as Z-type ligands or pendant groups are included in the relevant elementary reaction sections. While not exhaustive, the overall goal of this article is to provide a concise summary of the impressive variety of OM reactions that can be facilitated and enabled by LAs, with representative examples of each.

1.18.2

Oxidative addition

LAs increase the rate of OA of a large number of substrates to OM complexes. This enhancement is attributed to the withdrawal of electron density from the substrate by the LA, which weakens the bond that is being broken and therefore increases the thermodynamic favorability of OA, which also typically results in faster kinetics. This effect has been exploited in a range of catalytic reactions, including hydrocyanation (Section 1.18.2.2), hydroarylation/hydroalkylation (Section 1.18.2.3), carbocyanation (Section 1.18.2.4), cross-coupling (Section 1.18.2.5), cycloadditions (Section 1.18.2.6), and carbonylation (Section 1.18.2.7).

1.18.2.1

Stoichiometric oxidative addition

Jacobi and Ghosh investigated the OA of thioimidates to Pd(0) promoted by Lewis acidic alkyl Zn(II) iodides.10 The thioimidates studied did not undergo OA in the absence of a LA, and the E-isomers of thioimidates did not react even in the presence of a LA. In contrast, stoichiometric EtZnI promotes the OA of the CdS bond in the Z-isomers of thioimidates to Pd(PPh3)4. X-ray diffraction and nuclear Overhauser effect data indicate that the Z-isomer is able to form an equilibrium mixture, with the thioimidate bound to the Zn in two different bidentate conformations (Scheme 2A). Specifically, along with binding of the thioimidate, a pyrrole ring appended to the thioimidate coordinates in either an Z1 or Z5 fashion. The bidentate interaction between the LA and the thioimidate is proposed to facilitate OA by increasing the polarization of the CdS bond. As the E-isomer is sterically unable to bind to Zn in a bidentate fashion, it does not undergo OA (Scheme 2B).

(A)

(B)

Scheme 2 Lewis acid/base adducts formed by (A) Z and (B) E-thioimidates. In the case of Z-thioimidates the ability to coordinate in a bidentate fashion facilitates OA to Pd.

558

Lewis Acid Participation in Organometallic Chemistry

While studying cross-coupling reactions involving thiopyridines, Jacobi and Ghosh observed that OA of 2- and 3-substituted pyridines to Pd(0) only occurs in the presence of a LA such as BF3 ∙Et2O (Scheme 3).10 Coordination of the LA to the pyridyl nitrogen decreases the electron richness of the substrate and polarizes the CdS bond sufficiently to enable OA. In related work, Ogoshi and coworkers11 studied the OA of CdF bonds to Ni and Pd bisphosphine complexes, and observed that the C(sp2)dF bond of tetrafluoroethylene (TFE) is cleaved under mild conditions in the presence of various LAs. Specifically, the presence of 1–2 equivalents of MgBr2, AlCl3, LiI, BF3 ∙ Et2O, and B(C6F5)3 all enable OA at room temperature or with heating to 40  C, while the same reaction without a LA requires heating to 100  C (Scheme 4). Overall, these studies demonstrate the dramatic effects that LAs can have in facilitating OA reactions.

Scheme 3 OA of 2- and 3-substituted pyridines to Pd(0) enabled by BF3.

Scheme 4 OA of CdF bonds facilitated by LAs.

1.18.2.2

Hydrocyanation

LAs are proposed to enhance the rate of both OA and RE in catalytic hydrocyanation reactions (see also Section 1.18.3.2). For example, Morandi et al. reported the reversible transfer hydrocyanation of alkenes using a Ni(0) catalyst and AlMe2Cl as a LA co-catalyst (Scheme 5A).12 Careful selection of the HCN donor or alkene is required to drive the equilibrium reaction and achieve high yields, but the method is valuable because it avoids the use of toxic HCN gas by transferring the equivalent of HCN from a donor molecule. A subsequent report demonstrated that a slightly modified catalytic system is active for Mizoroki-Heck reactions of aryl (Ar) cyanides (Scheme 5B).13 Here, the LA is proposed to be vital for OA of the CdCN bond. Specifically, on the basis of density functional theory (DFT) calculations Dang et al. suggested that the LA coordinates to the CN nitrogen atom (Scheme 6).14 This interaction lowers the barrier for CdCN OA by 19.6 kcal/mol compared to the LA-free case. An additional example of catalytic transfer hydrocyanation of alkenes and alkynes was reported by the Studer group with a Pd(0) catalyst, Lewis acidic BPh3 as a co-catalyst, and 1-methylcyclohexa-2,5-diene-1-carbonitrile as an HCN source (Scheme 5C).15 The LA is again proposed to coordinate to the CN nitrogen atom and enhance the rate of OA (Scheme 6). These examples demonstrate the utility of LAs in hydrocyanation catalysis. (A)

(B)

(C)

Scheme 5 (A) Reversible transfer hydrocyanation of alkenes. (B) Aryl cyanides as electrophiles in Mizoroki-Heck reactions. (C) Transfer hydrocyanation with 1-methylcyclohexa-2,5-diene-1-carbonitrile.

Lewis Acid Participation in Organometallic Chemistry

559

Scheme 6 Coordination of LAs to promote OA of nitrile substrates.

1.18.2.3

Hydroarylation and hydroalkylation

There are many examples of the use of LAs as co-catalysts in hydroalkylation and hydroarylation reactions (Scheme 7A).16 In 2008, the groups of Hiyama and Nakao utilized a phosphine-supported Ni(0) catalyst for the C-2 alkenylation of pyridines in the presence of AlMe3, ZnMe2, or ZnPh2 as co-catalysts.16a The LA is proposed to coordinate to the pyridine nitrogen atom, withdrawing electron density from the substrate and promoting OA of the proximal CdH bond (Scheme 7B). Modification of this system allowed for selective C-4 alkylation of pyridines through the use of the bulky LA methylaluminum bis(2,6-di-tert-butyl-4-methylphenoxide) (MAD), and a Ni catalyst with an N-heterocyclic carbene (NHC) ligand (Scheme 7B).16c In related chemistry, the alkenylation and alkylation of 2-pyridones was achieved.16b In this case, the AlMe3 is proposed to coordinate to the carbonyl oxygen atom, which promotes OA of the proximal C(6)dH bond (Scheme 7C). The selective C-4 alkylation of benzamides and aromatic ketones was also achieved using Ni(NHC) catalysts and MAD.16e DFT studies support that coordination of the LA to the carbonyl oxygen accelerates OA of the C(4)dH bond (Scheme 7D).

(A)

(B)

(C)

(D)

Scheme 7 (A) Generic example of a hydroalkylation or hydroarylation reaction involving a LA co-catalyst. (B) LA-dependent C-2 or C-4 OA of pyridines. (C) OA of pyridones at C-6. (D) OA of benzamides/ketones at C-4.

1.18.2.4

Carbocyanation

Hiyama, Nakao, and co-workers initially reported the catalytic carbocyanation of alkynes and alkenes using a Ni(0) catalyst and a LA co-catalyst, such as BPh3, AlMe3, or AlMe2Cl (Scheme 8).17 By changing the transition metal (TM) catalyst, LA, and reaction conditions, the substrate scope of these reactions was greatly expanded and now includes both inter- and intramolecular substrates.18 In all cases, the LA is proposed to coordinate to the CN nitrogen atom and facilitate OA, analogous to hydrocyanation (Scheme 6). The groups of Douglas and Morandi further expanded the scope of these reactions to include intramolecular cyanoamidation19 and aminocyanation20 of alkenes and cyanation of aryl chlorides with butyronitrile to generate aryl nitriles.21 The Guan group explored the mechanism for the cyanoesterification of alkynes (Scheme 8, R ¼ ester) via DFT studies.22 The presence of a LA such as BAr3, which coordinates to the CN group throughout the reaction, changes the nature of the turnover-limiting step and lowers the overall barrier for the process. However, contrary to other examples in this section, the barrier for OA of the CdCN bond is higher (25.4 kcal/mol) in the presence of a LA than in the absence of a LA (21.2 kcal/mol). This is

560

Lewis Acid Participation in Organometallic Chemistry

compensated by the fact that the barrier to migratory insertion (now the turnover-limiting step) decreases from 31.0 kcal/mol in the absence of a LA to 16.7 kcal/mol in the presence of a LA. This example demonstrates that it is not universally true that a LA lowers the barrier to OA, and that even if a LA raises the barrier to one elementary step it can still improve the rate of the overall catalytic reaction. Subsequently, the same group examined the effect of LAs on the polyfluoroarylcyanation of alkynes (Scheme 8, R ¼ C6F5).23 In this case, DFT studies indicate the LA significantly decreases the barrier to OA of the CdCN bond: 17.4 vs. 22.0 kcal/mol without a LA.

Scheme 8 Representative substrates used in catalytic carbocyanation.

1.18.2.5

Cross-coupling

In an analogous fashion to hydrocyanation reactions, LAs are proposed to enhance the rate of OA of the electrophile and RE of the product in cross-coupling reactions (Section 1.18.3.4). For example, the Martin group reported Kumada couplings involving arylation of 2,3-dihydrofuran CdO bonds using a Ni(0) catalyst and stoichiometric LiCl (Scheme 9A).24 Mechanistic studies indicate that Li+ assists in OA of the CdO bond through coordination to the oxygen atom. In related work, Shu and co-workers described cross-electrophile coupling reactions of Ar halides and allylic alcohols in which a ZrCl4 co-catalyst activates the alcohol for OA of the CdO bond by the Ni(0) catalyst.25 Without a LA, undesired biaryl byproduct is exclusively observed (Scheme 9B).

(A)

(B)

Scheme 9 (A) Kumada coupling of 2,3-dihydrofuran in the presence of a LA. (B) Cross-electrophile coupling of Ar halides and allylic alcohols with and without a LA.

1.18.2.6

Cycloaddition

LA co-catalysts can promote TM-catalyzed cycloaddition reactions. For example, Hiyama and Nakao et al. described the dehydrogenative [4 + 2] cycloaddition of formamides and alkynes using a Ni(0) catalyst and AlMe3 (Scheme 10A).26 The LA is proposed to

Lewis Acid Participation in Organometallic Chemistry

561

coordinate to the oxygen of the formamide carbonyl, promoting OA of the formyl CdH bond (Scheme 10B). A DFT study on this reaction showed that coordination of the LA to the formamide lowers the barrier for CdH OA by 14 kcal/mol compared to the LA-free reaction.27 Subsequently, the Ogoshi group proposed that AlMe2Cl plays a similar role in the [3 + 2] cycloaddition of cyclopropyl ketones and alkynes (Scheme 10C).28 Specifically, the AlMe2Cl coordinates to the carbonyl oxygen atom of the cyclopropyl ketone, thereby facilitating OA of the cyclopropyl CdC bond.

(A)

(B)

(C)

(D)

Scheme 10 (A) [4 + 2] Cycloaddition of formamides and alkynes. (B) OA of a formyl CdH bond to Ni(0). (C) [3 + 2] cycloaddition of cyclopropyl ketones and alkynes. (D) [4 + 2] cycloaddition of o-cyanophenylbenzamide derivatives with alkynes.

A study by Kurahashi and Matsubara et al. described the use of MAD as a LA co-catalyst to promote the formation of quinolones via the [4 + 2] cycloaddition of o-cyanophenylbenzamide derivatives with alkynes through loss of Ar-CN as a byproduct (Scheme 10D).29 A complementary DFT study proposed that one equivalent of the LA coordinates to the CN nitrogen atom and promotes C–C OA of the substrate.30 A second equivalent of the LA coordinates to the carbonyl oxygen atom, promoting C–C coupling and b-Ar elimination at the metal.

1.18.2.7

Carbonylation

Alper et al. reported the catalytic carbonylation of epoxides using a catalyst formed in situ from Co2(CO)8 and bis(triphenylphosphine)iminium, ([Ph3P¼]2N, PPN).31 Specifically, when the Co system is used in combination with BF3 ∙ Et2O as a LA co-catalyst, a variety of b-lactones were synthesized with high regioselectivity from monosubstituted epoxides and CO (Scheme 11A). Similarly, the Coates group reported the use of [LA][Co(CO)4] catalysts for epoxide carbonylation and found that the optimal LA is an Al(salen) (salen ¼ N,N0 -ethylenebis(salicylimine)) complex.32 A mechanistic study from Molnar, Luinstra, Rieger et al. found that strong LAs increase the rate of epoxide ring opening, but decrease the rate of lactone ring closing (Scheme 11B).33 Conversely, weak LAs hinder epoxide ring opening and promote facile lactone ring closing. In agreement with this proposal, the Coates group published a mechanistic study on their [Al(salen)][Co(CO)4] system and determined that epoxide opening is facile with the strong LA, resulting in lactone ring closing being turnover-limiting.34 In related chemistry, Li et al. demonstrated catalytic carbonylation of NMe3 with [Co(CO)4]−, Me4NI, and a variety of co-catalytic LAs including AlCl3, FeCl3, BiCl3, and InCl3 (Scheme 11C).35 All LAs studied improved the selectivity for N,N-dimethylacetamide (DMAc) over N,N-dimethylformamide (DMF) or acetic acid, but AlCl3 was found to afford a slightly higher yield of DMAc. The LA is proposed to coordinate to the nitrogen atom of the amine, promoting OA of the CdN bond.

562

Lewis Acid Participation in Organometallic Chemistry

(A)

(B)

(C)

Scheme 11 (A) LA-assisted carbonylation of epoxides. (B) LA-dependent turnover-limiting steps: OA is turnover-limiting with weak LAs, RE is turnover-limiting with strong LAs. (C) Selective carbonylation of trimethylamine to DMAc.

1.18.3

Reductive elimination

In a similar manner to OA, the rates of the microscopic reverse reaction, RE, can be increased through the use of LAs. The LA pulls electron density away from the ligands and metal center, subsequently lowering the barrier for RE. This effect has been exploited catalytically in hydrocyanation (Section 1.18.3.2), decarbonylation (Section 1.18.3.3), and cross-coupling (Section 1.18.3.4) reactions. It is also briefly discussed in Section 1.18.2.7 on carbonylation reactions.

1.18.3.1

Stoichiometric reductive elimination

The Bercaw group studied the effects of exogenous Brønsted and Lewis acids on RE from a hydridoalkylplatinum(IV) intermediate relevant to alkane oxidation.36 They discovered that not only do protic acids such as HCl and triflic acid promote RE, but the LA SnCl4 also dramatically increases the rate of RE (Scheme 12). In work related to olefin hydrocyanation (see Section 1.18.3.2), Nolan and coworkers investigated the effects of LAs on the rate of RE from cyanoalkylpalladium(II) complexes (Scheme 13).8a They studied a range of LAs: BEt3, BPh3, AlEt3, B(C6F5)3, and AlPh3 (all abbreviated ER3), and a linear correlation was observed between the rate of RE from (dppp)Pd(CH2SiMe3)(CNER3) (dppp ¼ 1,3-bis(diphenylphosphino)propane) and the Lewis acidity of ER3. In this case, Lewis acidity was quantified by determining the enthalpy of LA binding to the cyanide complex. AlPh3, the strongest LA studied, increased the rate of RE 60-fold compared to the additive-free reaction. Coordination of a LA is known to induce positive charge generation at nitrile ligands while creating negative charge buildup at the LA.37 Nolan et al. propose that, after LA coordination, the electrophilic carbon undergoes nucleophilic attack by the alkyl ligand to generate the alkylcyanide product, which then dissociates from the metal center.

Scheme 12 RE from a Pt(IV) complex facilitated by Brønsted acids or a LA.

Lewis Acid Participation in Organometallic Chemistry

563

Scheme 13 Proposed mechanism of LA-promoted RE of alkyl cyanides from Pd.

While seeking to improve cross-coupling reactions between amines and N-heterocycles, the Hartwig group explored the effect of LAs on the rates and yield of RE from pyridylpalladium amido complexes (Scheme 14).38 They observed that RE from 3- and 4-pyridylpalladium amido complexes is accelerated by boron LAs, with a greater enhancement for 4-pyridyl substrates. A positive correlation between Lewis acidity and the enhancement in rate was found, similar to Nolan’s work.8a Hartwig and Shen propose that the acceleration arises from LA binding to the pyridyl nitrogen, which results in a more electron-deficient Ar group that undergoes RE faster than more electron-rich systems lacking a LA.39

Scheme 14 RE from Pd(II) complexes accelerated by boron LAs.

Tilley and coworkers published two studies on the effect of LAs binding to 2,20 -bipyrazine (bpyz) or 2,20 -bipyrimidine (bpym) Pt biaryl complexes and acting as “remote triggers” for RE (Scheme 15).40 Their studies provide insight into electronic effects on the rates of RE in cross-coupling reactions. They observed that binding two equivalents of B(C6F5)3 to a (bpyz)Pt biaryl complex enables RE to occur at 45  C, while RE from the same Pt complex without a LA present requires heating above 100  C (Scheme 15A).40a Eyring analysis indicated a rate increase of 64,000 due to borane binding. Similarly, Tilley et al. investigated the influence on RE of several Si- and Zn-based LAs bound to (bpym)Pt biaryl complexes (Scheme 15B).40b In this case, LA binding increased the rate of RE by up to a factor of 108! The magnitude of the rate increase correlates with Lewis acidity. In related work, the Tilley and Bergman groups also observed that a Zn(II) LA acting as a Z-type ligand bound directly to a (biaryl)Pt(phen) (phen ¼ 1,10-phenanthroline) complex strongly accelerates the rate of biaryl RE, with complete conversion achieved in 15 min at 60  C (Scheme 16).41 Without a LA, no RE is observed over 48 h at 200  C. These results highlight that utilization of a bound or exogenous LA clearly can have a dramatic effect on the rate of RE (as well as its microscopic reverse OA, see Section 1.18.2.1), a ubiquitous and often turnover-limiting step in catalysis.

(A)

(B)

Scheme 15 LAs studied by the Tilley group act as remote triggers for increasing the rate of RE from Pt with (A) bpym and (B) bpyz ligands.

564

Lewis Acid Participation in Organometallic Chemistry

Scheme 16 Acceleration of biaryl RE from Pt(II) by a Zn LA.

1.18.3.2

Hydrocyanation

One of the earliest commercial uses of LAs in OM chemistry was to assist in controlling the selectivity in Ni-catalyzed hydrocyanation reactions of terminal alkenes (Scheme 17A).42 LAs such as ZnCl2 or AlCl3 were observed to affect the selectivity for terminal or internal alkene products by modifying the relative rate of alkene isomerization under the catalytic conditions. Further work by Tolman et al. showed LA additives also improved rates and catalyst lifetimes.43 Coordination of the LA to the nitrogen of the CN group on a hydrido Ni cyano intermediate was observed by IR and NMR spectroscopies, and is proposed to aid in RE. A report in 2009 from Pidko, Vogt et al. examined the hydrocyanation of styrenes with AlCl3 as a LA (Scheme 17B).44 Upon addition of LA, the linear product is exclusively observed, in contrast to the branched product, which is observed without LAs. DFT studies support that the pathway to linear product is both kinetically and thermodynamically favored in the presence of a LA. Overall, LAs can facilitate both RE and OA (see Section 1.18.2.2) in catalytic hydrocyanation reactions, allowing for milder conditions and selective formation of linear products.

(A)

(B)

Scheme 17 Linear hydrocyanation of (A) alkenes and (B) styrenes.

1.18.3.3

Decarbonylation

Kurahashi and Matsubara et al. described the catalytic decarbonylative addition of anhydrides to alkynes utilizing a Ni(0) catalyst and Zn(II) salts as LA co-catalysts (Scheme 18).45 Mechanistic studies suggest the most likely role of the LA is coordination to the carbonyl oxygen of the metallacycle, which withdraws electron density from the metal center through a conjugated system and promotes RE to the a-pyrone product.

Scheme 18 Catalytic decarbonylative addition of anhydrides and alkynes via LA-assisted RE from metallacycle to generate a-pyrone products.

1.18.3.4

Cross-coupling

LA co-catalysts, such as BEt3, enhance the rate of product formation in the amidation of heteroaromatic Ar halides using Pd catalysts (Scheme 19).38 Mechanistic studies indicate that coordination of a boron LA to the nitrogen of a pyridylpalladium intermediate enhances the rate of RE by three orders of magnitude. As a consequence, catalysis with a LA results in significantly higher yields of

Lewis Acid Participation in Organometallic Chemistry

565

product than the reaction without a LA; a representative substrate is shown in Scheme 19. A similar effect was observed by Knochel et al. in the catalytic cross-coupling of pyridines with Ar bromides, wherein a LA, such as ZnCl2, Sc(OTf )3 (OTf ¼ trifluoromethanesulfonate), or BF3 ∙OEt2, coordinates to the nitrogen of the pyridylpalladium intermediate.46 The coordination of the LA to the metal pyridyl species in all of these examples decreases the electron density at the metal center and promotes RE.

Scheme 19 Amidation of Ar halides via pyridylpalladium species.

1.18.4

1,1-Insertion

There are numerous examples of LA additives accelerating the rate of 1,1-insertion reactions, specifically those involving CO. Generally, the LA is proposed to bind to the nucleophilic oxygen of CO to pull electron density away from the metal center and facilitate migration of the alkyl group. These reactions have primarily been studied in stoichiometric reactions (Section 1.18.4.1.1), although LAs can improve catalytic systems for CO hydrogenation (Section 1.18.4.1.2), a reaction which is relevant to industrial syngas conversion.47

1.18.4.1 1.18.4.1.1

CO insertion Stoichiometric CO insertion

In 1972, the Collman group demonstrated that alkali cations increase the rate of CO insertion into Fe-alkyl bonds.48 They proposed that the Lewis acidic cation stabilizes negative charge buildup on the carbonyl oxygen that occurs during the insertion, thus stabilizing the transition state (TS) (Scheme 20). In subsequent years, detailed mechanistic studies provided support for this hypothesis.49 For example, the Shriver group reported the crystal structure of a LA, AlBr3, bound to the oxygen of the carbonyl ligand oxygen in (CO)4MnC(O)CH3.50

Scheme 20 CO insertion into Fe–alkyl bonds facilitated by alkali cations.

The rate of CO insertion can also be accelerated by LAs that are covalently bound to a ligand already coordinated to the metal center. An example containing a pendant LA was reported by Labinger, Miller, and coworkers in 1982.51 They synthesized a series of aluminoaminophosphines as amphoteric ligands for Fe and Mn carbonyl complexes, and found that the ligands enabled facile CO insertion into FedMe and MndMe bonds (Scheme 21A). The Lewis acidic Al moiety increases the rate of CO insertion, and the Lewis basic phosphine group fills the vacant coordination site generated during the reaction. More recently, the Bercaw group developed a variety of systems containing pendant LAs to accelerate CO insertion reactions on Re.52 They reported that a Lewis acidic borane in the secondary coordination sphere of a Re carbonyl complex facilitates spontaneous CO insertion into an Re–alkyl bond after the delivery of two hydride equivalents from either NaHBEt3 or [HPt(dmpe)2][PF6] (dmpe ¼ 1,2-bis(dimethylphosphino)ethane) (Scheme 21B).52a Although these examples have not resulted in the catalytic reductive coupling of CO, they represent major steps toward achieving Fischer-Tropsch chemistry under mild conditions and demonstrate the potential power of LAs to facilitate this chemistry.

566

Lewis Acid Participation in Organometallic Chemistry

(A)

(B)

Scheme 21 (A) CO insertion into an FedC bond enabled by an amphoteric aluminoaminophosphine ligand. (B) CO insertion into a RedC bond enabled by pendant borane LAs.

1.18.4.1.2

CO hydrogenation

Two consecutive reports from the group of Muetterties in 1981 discussed the catalytic hydrogenation of CO using LAs.53 The first system consisted of an Ir4(CO)12 catalyst dissolved in a NaCl/AlCl3 melt. Products ranged from methane (C1) to dimethylcyclohexane (C8), with the major product, isobutane (C4), produced in 54% yield. Without a LA, methane and ethane are generated but at a much slower rate. IR studies of the catalytic mixture indicate the presence of a CO stretch consistent with an Ir–CO–Al adduct (Scheme 22). It is proposed that this interaction enhances the 1,1-insertion step of the reaction. The second system consisted of an Os3(CO)12 catalyst in liquid BBr3. The products range from methane (C1) to neopentane (C5), as well as monobrominated hydrocarbons including methyl bromide and ethyl bromide. This system is proposed to involve the same type of CO-LA interaction.

Scheme 22 LA interaction with a metal carbonyl.

1.18.5

1,2-Insertion

LAs can drastically accelerate the rate of 1,2-insertion reactions. The electrophile is proposed to stabilize the negative charge buildup in the rate-determining TS for the insertion. This has been extensively studied for stoichiometric CO2 insertion reactions (Section 1.18.5.1), as well as the insertion of esters and carboxylic acids (Section 1.18.5.2). Catalytically, this effect has been exploited in a variety of transformations, including CO2 reduction (Section 1.18.5.1.2), formic acid and methanol dehydrogenation (Section 1.18.5.1.3), ester and carboxylic acid hydrogenation (Section 1.18.5.2.1), and amide hydrogenation (Section 1.18.5.2.2).

1.18.5.1 1.18.5.1.1

CO2 insertion Stoichiometric CO2 insertion

Stoichiometric CO2 insertion reactions into TM hydride, alkyl, and allyl bonds can be accelerated by the addition of a LA. Regardless of whether a LA is present, these reactions are proposed to proceed through a two-step mechanism, which is illustrated in Scheme 23A.54 When the TS associated with the first step is rate-determining, the insertion reaction is referred to as an outer-sphere process because the CO2 does not interact directly with the metal center (Scheme 23B). If the TS associated with the second step is rate-determining, the insertion reaction is referred to as an inner-sphere process because the CO2 interacts directly with the TM. Several detailed mechanistic studies have been performed to probe how LAs impact the rates of CO2 insertion for both inner- and outer-sphere procecess.54,55 In general, it is proposed that LAs only promote CO2 insertion for reactions that follow an outer-sphere pathway. Specifically, it was demonstrated that LAs do not change the rate of CO2 insertion into pincer-supported Ni complexes, which react via an inner-sphere pathway (Scheme 24A). In contrast, large increases in insertion rates are observed for the complexes (iPrPNHP)IrH3 (iPrPNHP ¼ HN(CH2CH2PiPr2)2) and (terpy)(bpy)RuH+ (terpy ¼ 2,20 :60 ,200 -terpyridine) (bpy ¼ 2,20 -bipyridine) (Scheme 24B and C), which insert via an outer-sphere pathway. In general, smaller alkali metal cations such as Li+ give more significant increases in insertion rates than larger cations such as Na+ and K+. Additionally, the effect of the LA on the rates of insertion decreases in more polar solvents. Thus, although the presence of Li+ gives a substantial increase in the rate of CO2 insertion into (terpy)(bpy)RuH+ in acetonitrile, there is almost no increase in either MeOH or DMF. These reports provide experimental guidance for determining whether or not a CO2 insertion reaction will be assisted by a LA, as well as guidance for using LAs to improve the efficiency of catalytic CO2 reduction reactions.

Lewis Acid Participation in Organometallic Chemistry

567

(A)

(B)

Scheme 23 (A) Mechanism of CO2 insertion into metal-element s-bonds. (B) Stabilization of the outer-sphere rate-determining TS by a LA.

(A)

(B)

(C)

Scheme 24 (A) LAs do not enhance the rate of CO2 insertion into pincer-supported Ni complexes, which follow an inner-sphere pathway. LAs drastically increase CO2 insertion rates into (B) an Ir and (C) a Ru complex, which follow outer-sphere pathways.

1.18.5.1.2

CO2 reduction

LAs can enhance the catalytic reduction of CO2 in both electrochemical and thermal reactions.56 However, given the complexity of electrocatalytic systems, it is typically unclear how LAs present as part of the electrolyte aid electrochemical CO2 reduction, so these systems will not be discussed further. An example of a thermal reaction that is impacted by the presence of a LA is CO2 hydroboration.57 Although systems for hydroboration always contain a Lewis acidic borane, in most reactions the role of the borane is to drive the reaction by trapping thermodynamically unstable intermediates. Nevertheless, the identity of the borane directly impacts the product selectivity in CO2 hydroboration, as sterically bulky boranes tend to generate products in the same oxidation state as formic acid, whereas smaller

568

Lewis Acid Participation in Organometallic Chemistry

boranes often generate products in the same oxidation state as methanol (Scheme 25A).55b,58 Additionally, the use of a LA co-catalyst can alter the selectivity of CO2 hydroboration when a pincer-supported Ni catalyst is used in conjunction with 9-BBN (borabicyclo[3.3.1]nonane) as a reductant (Scheme 26).58c Specifically, 10 mol% B(OPh)3 changes the selectivity from a product in the same oxidation state as formaldehyde to a product in the same oxidation state as methanol. In this case, the B(OPh)3 LA is proposed to change the rates of 1,2-insertion of key catalytic intermediates.

(A) (B)

Scheme 25 Various products of CO2 reduction via (A) hydroboration and (B) hydrosilylation.

Scheme 26 LA promoted change in selectivity in catalytic CO2 hydroboration.

Hydrosilylation is another method for reducing CO2 under relatively mild conditions (Scheme 25B).59 Turculet et al. demonstrated that Pd and Pt complexes supported by pincer ligands can reduce CO2 to CH4 in the presence of an excess of a tertiary silane.60 Catalytic reduction is only observed in the presence of the LA B(C6F5)3, and the use of a weaker LA such as BPh3 did not generate reduced products. The LA associates with the metal hydride to generate an ion pair, and then coordinates to the carbonyl of the formate ligand to make CO2 insertion more thermodynamically favorable (Scheme 27A). The Berke group reported a Re catalyst with a B(C6F5) co-catalyst for the hydrosilylation of CO2 to bis(silyl) acetal.61 NMR studies support the formation of a Lewis pair, wherein the LA is coordinated to an oxygen atom of CO2 and promotes insertion into the metal hydride (Scheme 27B).

(A)

(B)

Scheme 27 (A) CO2 insertion into a metal hydride bound to a LA. (B) CO2 insertion into a metal hydride can also be facilitated through the binding of a LA to coordinated CO2.

1.18.5.1.3

Formic acid or methanol dehydrogenation

LAs can act as co-catalysts in both TM-catalyzed formic acid and methanol dehydrogenations. These reactions typically require the use of exogenous base or ligand, however, LA co-catalysts can also be used as additives. Under optimized conditions, the pincer-supported Fe complex, A, and a LiBF4 co-catalyst achieved a turnover frequency (TOF) of almost 200,000 and a turnover number (TON) of nearly 1,000,000 for formic acid dehydrogenation (Scheme 28A).62 Mechanistic studies indicate the LA aids in turnover-limiting decarboxylation from A, the microscopic reverse of CO2 insertion (Scheme 28B).

Lewis Acid Participation in Organometallic Chemistry

569

(A)

(B)

Scheme 28 (A) Catalytic formic acid dehydrogenation with A. (B) Decarboxylation of A via a H-bound formate intermediate.

LAs can also act as co-catalysts in aqueous methanol dehydrogenation (Scheme 29).63 Improved rates and a total TON of 51,000 were achieved using A in conjunction with a LiBF4 co-catalyst. The LA is again implicated in facilitating decarboxylation of the Fe formate. DFT studies indicate that the LA stabilizes the negative charge buildup in an H-bound formate intermediate (Scheme 28B). In related work, the Beller group described a pincer-supported Ru catalyst, B, that operates in 8 M KOH.64 Although no role was initially proposed for the K+, a subsequent computational study suggested that KOH has a dual role; the OH− acts as a base to deprotonate the substrate and the K+ acts as a LA to enable H2 formation (see Section 1.18.9.2).65

Scheme 29 Catalytic methanol dehydrogenation with A or B.

1.18.5.2 1.18.5.2.1

Carbonyl insertion Ester or carboxylic acid hydrogenation

The Beller group reported that a Ru triphos system (triphos ¼ bis(diphenylphosphinoethyl)methylphosphine), in conjunction with an Al(OTf )3 co-catalyst, is active for the hydrogenation of esters to form ethers with water as a byproduct (Scheme 30A).66 Mechanistic studies indicate that the LA is crucial for selective CdOH bond cleavage, as it binds to the carbonyl oxygen atom of the ester, which increases the rate of ester insertion into the RudH bond to form a hemiacetal intermediate (Scheme 31). Further reduction of the hemiacetal is likely also aided by the LA. The Beller group extended this system to carboxylic acids, using Sn(OTf )2 in place of Al(OTf )3 to generate primary alcohols (Scheme 30B).67 This LA is proposed to coordinate to the carbonyl moiety in an analogous fashion.

(A)

(B)

Scheme 30 Hydrogenation of (A) esters to ethers and (B) carboxylic acids to alcohols.

570

Lewis Acid Participation in Organometallic Chemistry

Scheme 31 Hydrogenation of esters or carboxylic acids (R2 ¼ H) via a hemiacetal intermediate to ethers or alcohols, respectively.

1.18.5.2.2

Amide hydrogenation to amines

The hydrogenation of amides can proceed through either a pathway involving CdO or CdN bond cleavage (Scheme 32A). The Beller group described the hydrogenation of amides to primary or secondary amines in the presence of LAs via C–O cleavage.68 Specifically, Yb(OTf )3 and a triphos-supported Ru catalyst react under milder reaction conditions to give improved yields and selectivities compared to reactions only using Ru and triphos (Scheme 32B). Although the LA is essential for the reaction, its role is unclear. Subsequently, the group of Zhou utilized the same Ru catalyst with a BF3 ∙Et2O co-catalyst for the hydrogenation of amides to secondary amines.69 Mechanistic studies indicate an interaction between the carbonyl group of the amide and the LA, and hydrogenation of this adduct is highly selective to the secondary amine via C–O cleavage rather than the primary amine via C–N cleavage. In 2018, Milstein et al. reported a pincer Mn system for the hydrogenation of amides to secondary amines via C–O cleavage with co-catalytic tBuOK and B(C6F5)3 (Scheme 32C).70 The Lewis acidic B(C6F5)3 is proposed to coordinate to the oxygen atom of the amide, enhancing the 1,2-insertion of the carbonyl into the metal hydride (Scheme 33).

(A)

(B)

(C)

Scheme 32 (A) Amide hydrogenation via C–O cleavage or C–N cleavage. (B) Ru-catalyzed amide hydrogenation to amines. (C) Mn-catalyzed amide hydrogenation to amines.

Scheme 33 LA-assisted 1,2-insertion of an amide into Mn hydride to form an hemiaminal.

1.18.6

Ligand addition or abstraction

LAs are known to assist in ligand addition and abstraction reactions. The electrophile is proposed to withdraw electron density from the reacting ligand, which weakens the metal-ligand bond and results in a new LA-ligand bond. In catalytic reactions, ligand abstraction by LAs is frequently exploited in order to activate OM precatalysts by creating an open site for substrate binding. This elementary reaction has been utilized in catalytic olefin polymerization (Section 1.18.6.1.2), olefin trimerization (Section 1.18.6.1.3), and amide arylation (Section 1.18.6.2.2) reactions. Ligand abstraction reactions are closely related to the ligand substitution reactions presented in Section 1.18.7. The difference is that in the ligand abstraction reactions described in this section an obvious coordinating ligand is not added to the reaction mixture, while in the ligand substitution reactions presented in Section 1.18.7, a well-defined ligand is added to the reaction mixture in addition to the LA.

Lewis Acid Participation in Organometallic Chemistry

1.18.6.1 1.18.6.1.1

571

Alkyl abstraction Stoichiometric alkyl abstraction

Exogenous LAs can abstract alkyl groups from OM complexes and generate an open coordination site. The Goldberg group used B(C6F5)3 to abstract a methyl (CH−3) ligand from a Pt(II) complex to facilitate the OA of arenes and alkanes to generate Pt(IV) alkyl hydrides (Scheme 34).71 Similarly, the Marks group used B(C6F5)3 to promote methyl abstraction from Zr and Hf alkyl metallocene complexes relevant to olefin polymerization (see Section 1.18.6.1.2).72 They observed that abstraction rates depend on both the solvent dielectric constant and the identity of the alkyl group. Alkyl group abstraction has also been reported using LAs tethered to the metal center.73

Scheme 34 Methyl abstraction by B(C6F5)3 to enable OA of alkanes and arenes to Pt(II).

1.18.6.1.2

Olefin polymerization

In 1994, Marks and co-workers examined the effect of LA additives in olefin polymerization.74 They described a Zr precatalyst that is activated through methyl abstraction by B(C6F5)3 (Scheme 35A). Subsequent H2 addition yields a cationic Zr–H, which is highly active for both catalytic ethylene polymerization and atactic propylene polymerization. An additional report from the Marks group examined the effect of different Ar groups on the enthalpy change associated with the reaction of B(C6F5)2Ar with complexes of the form Cp0 2MMe2 (Cp0 ¼ 1,2-Me2Cp, Cp ; M ¼ Ti, Zr, Hf ) (Scheme 35B), as well as the catalytic activity of [(Me2Cp)2MMe]+[MeB(C6F5)2Ar]− complexes in ethylene polymerization.75 Catalytic activity correlates with the substitution of Ar, with more Lewis acidic species giving higher activities (Ar ¼ C6F5 > 3,5-C6H3F2 > Ph  3,5-C6H3-Me2). Additionally, Zr complexes are more active than Hf species and a positive correlation between polymerization activity and the enthalpy of methyl abstraction was identified. The Marks group also described the effect of using sterically bulky borane and aluminate LAs in catalytic olefin polymerization.76 Bulky LAs decrease the extent of ion pairing between the LA-methyl complex and the resultant cationic metallocene, and as a result give higher activity for polymerization compared to less sterically hindered systems.

(A)

(B)

Scheme 35 (A) Precatalyst activation by BAr3 and H2 for ethylene and propylene polymerization. (B) Formation of Lewis pairs by methyl abstraction.

1.18.6.1.3

Olefin trimerization

In 2013, the Bercaw group utilized Ti precatalysts of the form [(FI)TiMe2][MeB(C6F5)3] (FI ¼ phenoxy-imine), which are synthesized through methyl abstraction from a FI trimethyl species by LAs such as B(C6F5)3, for ethylene and a-olefin trimerization (Scheme 36A).77 These complexes are selective for both ethylene trimerization to 1-hexene, as well as the unusual trimerization of linear a-olefins. For example, the system catalyzes the trimerization of 1-pentene, 1-hexene, and 1-decene with >95% selectivity, and gives an 85% yield of just one regioisomer. Modified precatalysts, including [(FI)Ti(CH2SiMe3)2Me] and [(FI)Ti(CH2CMe3)2Me], were later reported for the trimerization of ethylene and 1-pentene (Scheme 36B).78 The neopentyl analogue generates a significantly more active catalyst upon activation with B(C6F5)3; however, no stable products were isolated from the reaction with B(C6F5)3 and a variety of inactive side products were identified. The more efficient initiation of the neopentyl complex is offset by its decreased stability in trimerization catalysis. Overall, the activation of these complexes via methyl abstraction by a LA generates a catalyst that is highly efficient and selective.

572

Lewis Acid Participation in Organometallic Chemistry

(A)

(B)

Scheme 36 (A) Trimerization of ethylene (R ¼ H) and a-olefins. (B) Modified catalyst system with bulky alkyl groups neopentyl (X ¼ C) or trimethylsilylmethyl (X ¼ Si).

1.18.6.2 1.18.6.2.1

Halogen abstraction or exchange Stoichiometric halogen abstraction from metal bound CdX bonds

Shriver and Richmond studied stoichiometric halogen abstraction from OM complexes using LAs.79 They observed that the reaction of BX3 (X ¼ Cl, Br, I) with perfluoroalkyl metal complexes led to regiospecific activation of the CdF bond a to the metal center, forming BX2F and a CdX bond (Scheme 37). Similarly, the Baker group studied fluoride abstraction from metallacyclic OM complexes using LAs.80 They found that trimethylsilyl triflate (Me3SiOTf ) abstracts a fluoride from fluorinated Co cyclobutanes to provide products that vary with the metallacyclobutane substitution (Scheme 38A).80a They then studied the reactivity of nickel perfluorocyclopentanes with Me3SiOTf and observed activation of the a CdF bond, analogous to Shriver and Richmond’s work (Scheme 38B).80b These stoichiometric studies demonstrate the power of LAs to selectively activate carbon-halogen bonds in OM complexes.

Scheme 37 An example of selective a-halogen abstraction by BCl3.

(A)

(B)

Scheme 38 (A) Selective fluoride abstraction from Co metallacyclobutanes by Me3SiOTf. (B) Selective fluoride abstraction from Ni metallacyclopentanes by Me3SiOTf.

Lewis Acid Participation in Organometallic Chemistry

1.18.6.2.2

573

Arylation of amides

Dobereiner et al. examined the use of a Pd catalyst with M(OTf )3 salts in catalytic amide arylation (Scheme 39A).81 Although product is obtained in the absence of LAs, an increased yield is generally observed in the presence of a LA (i.e., 58% vs. 99%). Mechanistic studies suggest that the LA aids in turnover-limiting halide abstraction from Pd, which facilitates the binding of the amide (Scheme 39B).

(A)

(B)

Scheme 39 (A) Pd-catalyzed arylation of amides in the presence of a LA. A representative substrate is shown. (B) Halide abstraction by a LA additive.

1.18.6.3 1.18.6.3.1

Hydride abstraction Stoichiometric hydride abstraction for hydricity determination

Cationic LAs are commonly used to abstract a hydride from OM complexes, which can be valuable for measuring thermodynamic and kinetic hydricities.82 Thermodynamic hydricity is the free energy required to release a hydride ion from a given species, and can be used to predict whether an OM complex will transfer a hydride to a substrate.83 Kinetic hydricity measures the rate of hydride transfer and is directly related to understanding the rates of many catalytic hydrogenation and dehydrogenation reactions in which hydride transfer is the turnover-limiting step. One frequently used LA for hydride abstraction is the trityl cation, which is particularly useful because its hydride accepting ability can be tuned by changing the substitution on the phenyl groups.84 For example, the kinetic and thermodynamic hydricities of a variety of OM complexes were determined by the Bruno group using unsubstituted and tris(p-methoxy)-substituted trityl cations (Scheme 40).85 More electron-donating ligands on the metal center facilitate the hydride abstraction, and the reaction is both kinetically and thermodynamically more favorable with PMe3 compared to PPh3. The insights gained from these studies were used to understand hydride transfer in the hydrogenation of alkenes. In general, hydricity determination via LA-assisted hydride abstraction is a powerful technique that can provide critical insight into catalytic reactions that involve hydride transfer.

Scheme 40 Kinetic and thermodynamic hydricity of a representative Mo complex determined using the trityl cation.

1.18.7

Ligand substitution

LAs can influence both the rate and thermodynamics of ligand substitution reactions. For example, in stoichiometric reactions, this phenomenon has been studied using a pincer ligand with a hemilabile pendant crown ether that facilitates ligand substitution upon cation binding (Section 1.18.7.1). The same system has also been used to catalyze olefin isomerization (Section 1.18.7.2) and

574

Lewis Acid Participation in Organometallic Chemistry

methanol carbonylation (Section 1.18.7.3). Conversely, the addition of a LA has been found to induce NO ligand bending, formally oxidizing the metal center and opening a coordination site to enable ligand substitution, which facilitates alkene hydrogenation (Section 1.18.7.4). LA additives also assist in the substitution of a formate ligand in catalytic CO2 hydrogenation (Section 1.18.7.5). The ligand substitution reactions in this section are closely related to the ligand abstraction reactions discussed previously (see Section 1.18.6) as the first step in the substitution reaction often involves the abstraction of a ligand by a LA.

1.18.7.1

Stoichiometric ligand substitution

The Miller group developed Ir and Ni pincer complexes with crown ether-containing ligands that undergo ligand substitution reactions upon binding of Lewis acidic cations to the crown ether.86 The concentration and identity of the added LA strongly affects the rate of ligand substitution and the corresponding change in crown ether coordination. For example, Miller et al. studied Ni pincer complexes containing hemilabile aza-crown ether macrocycles for LA-controlled reversible nitrile ligand substitution.86c They discovered drastically different cation binding affinities depending on the coordination mode of the crown ether, with the tridentate coordination mode binding LAs over 105 times more tightly than the tetradentate coordination mode. In noncoordinating solvents such as dichloromethane, addition of Li+ was found to thermodynamically stabilize nitrile ligand binding to Ni by at least 7 kcal/mol. This interaction was used to develop an in situ switchable pentafluorobenzonitrile binding/ release system through the addition of Li+/12-crown-4, as illustrated in Scheme 41. These examples illustrate the powerful effect that LAs can have on the coordination environment and ligand substitution rates of OM complexes.

Scheme 41 Reversible nitrile ligand binding/release enabled by the addition of Li+/12-crown-4.

1.18.7.2

Olefin isomerization

Pincer-crown ether Ir complexes are able to catalytically isomerize allylbenzene (Scheme 42).87 Addition of a co-catalytic amount of LiBArF4 (BAr-F4 ¼ tetrakis[(3,5-trifluoromethyl)phenyl]borate) provides up to a 1000-fold enhancement in the rate of isomerization and produces trans-b-methylstyrene in 99% selectivity. Comparatively, NaBArF 4 provides only a threefold rate enhancement, and KBArF 4 does not enhance the rate of isomerization. The LA coordinates to the crown ether moiety and opens a vacant site at the metal center for the alkene to bind. Addition of a chloride source can effectively switch “off” catalytic activity by precipitating LiCl or NaCl, and further addition of Li+ or Na+ switches catalytic activity back “on.”

Scheme 42 Olefin isomerization catalyzed by a pincer-crown ether Ir complex in the presence and absence of a LA co-catalyst.

Lewis Acid Participation in Organometallic Chemistry

1.18.7.3

575

Carbonylation

The Miller group also utilized pincer-crown ether Ir complexes to catalytically carbonylate methanol to acetic acid (Scheme 43).88 Mechanistic studies indicate that a LA co-catalyst, such as La(OTf )3, can play several roles: (i) it can coordinate to the crown ether, which stabilizes acyl products and weakens the Ir amine bond, and (ii) it can increase the rate of CO migratory insertion (see Section 1.18.4.1). Additionally, when the crown ether in the pincer ligand is changed from aza-15-crown-5 to aza-18-crown-6 the resulting Ir complex displays enhanced selectivity for methyl acetate over acetic acid. In this case, either co-catalytic LiCl or HfCl4 increase the total TON to 1190, with methyl acetate produced as the major product.

Scheme 43 Methanol carbonylation with pincer-crown ether Ir complexes.

1.18.7.4

Alkene hydrogenation

The Berke group examined a series of ReI2(NO)(PR3)2L (R ¼ iPr, Cy; L ¼ H2O, H2) complexes in which a LA causes bending of the nitrosyl ligand (Scheme 44A).89 This change from a formally NO+ to NO− ligand effectively oxidizes the Re center by two electrons and creates an additional free coordination site at Re. The Re complexes in conjunction with a LA co-catalyst were utilized to facilitate the hydrogenation of terminal alkenes such as 1-hexene, 1-octene, and styrene, as well as cyclic alkenes such as cyclohexene, cyclooctene, and 1,5-cyclooctadiene (Scheme 44B). The optimal LA is generated in situ by mixing Me2PhSiH and B(C6F5)3 to form Me2PhSi+ and HBðC6 F5 Þ3 − . Mechanistic studies indicate that coordination of the silylium cation induces nitrosyl bending and consequently H2 cleavage.89a A subsequent report utilized Re(H)(NO)2(PR3)2 (R ¼ iPr, Cy) complexes which undergo the same nitrosyl bend and examined a range of LAs including B(C6F5)3, [Et3O][B(C6F5)4], and Et3Si+, which was generated in situ from Et3SiH and B(C6F5)3.89b The highest TOFs for catalytic alkene hydrogenation are observed with [Et3O][B(C6F5)4], followed by Et3Si+, and then B(C6F5)3, which correlates with the calculated bond strength between the LA and the nitrosyl oxygen. Modification of the ancillary ligands around Re allows for the isolation of complexes in which the LA is bound to the NO ligand. Specifically, both B(C6F5)3 and Et+ (formed from [Et3O][B(C6F5)4]) adducts of [ReCl(PR3)2(NO)2)] (R ¼ iPr, Cy) were isolated and have bent NO ligands.89c

(A)

(B)

Scheme 44 (A) LA-induced nitrosyl bend. (B) Catalytic alkene hydrogenation with a Re complex.

Fan et al. also demonstrated that LA co-catalysts can promote TM-catalyzed hydrogenation reactions.90 Addition of NaBF4 to a Rh catalyst with a chiral aza-crown phosphoramidite ligand (Scheme 45A) led to association of the Na+ with the aza-crown moiety, which results in dissociation of the crown from the Rh. The interaction with Na+ is proposed to be crucial in the catalytic asymmetric hydrogenation of dehydroamino acid esters (Scheme 45B). Less than 1% conversion is observed in the absence of the Na+ bound to

576

Lewis Acid Participation in Organometallic Chemistry

the aza-crown compared to full conversion, with enantiomeric excesses up to 98%, in the presence of Na+. Addition of [2,2,2]-cryptand removes the Na+ from the catalyst, therefore “switching off” catalysis, and further addition of Na+ can turn catalysis back “on.”

(A)

(B)

Scheme 45 (A) Catalyst activation by addition of LA and deactivation by addition of cryptand to remove LA. (B) Asymmetric hydrogenation of dehydroamino acid esters.

1.18.7.5

CO2 hydrogenation

The direct hydrogenation of CO2 to formate or methanol can be promoted by the presence of LA co-catalysts. The Byers group used inorganic salts, such as KOAc, KHCO3, and K2CO3, in the hydrogenation of CO2 to formate with several different Ru and Fe catalysts.91 However, the proposed mechanism did not include a role for the cation and instead focused on the potential role of the anion. Several subsequent reports indicated that LA co-catalysts increase the TON and TOF of pincer-supported Fe and Co catalysts for the hydrogenation of CO2 to formate in the presence of excess 1,8-diazabicyclo[5.4.0]undec-7-ene (DBU) (Scheme 46).92 A variety of alkali metal LAs were examined, and LiOTf was found to give a TON of nearly 59,000 with a tertiary amine-containing pincer ligand on Fe (C) and 1500 for a pincer ligand containing a secondary amine (D). Without LA, the TON is less than 3000 for complex C. In both systems, the LA is proposed to bind to the carbonyl group in the Fe formate intermediate and promote formate release (Scheme 47). Two additional reports examined related Co complexes, E and F, for the hydrogenation of CO2 to formate in the presence of excess DBU and a LA co-catalyst.93 Again, the complex containing a tertiary amine, E, is a more productive catalyst with a TON of 29,000, while the complex with a pincer ligand containing a secondary amine, F, achieves a TON of only 450. The LA is proposed to act in analogous fashion to the Fe systems. Beller et al. have published a triphos Co system (G) with a LA co-catalyst, such as BF3•Et2O or Yb(OTf )3, which reduces CO2 to CH3OH.94

Scheme 46 Catalysts for CO2 hydrogenation to formate (C-F) and methanol (G).

Lewis Acid Participation in Organometallic Chemistry

577

Scheme 47 LA-induced formate release in catalytic CO2 hydrogenation.

1.18.8

b-Hydride elimination

Lewis acidic additives can increase the rate of b-hydride elimination from OM complexes. This is most commonly proposed to occur through the interaction of the electrophilic additive with an electron-rich chelating carboxylate ligand, which facilitates dechelation and opens the coordination site necessary for the elimination reaction to occur. In catalysis, this phenomenon can enable acrylate synthesis (Section 1.18.8.2) and allene carboxylation (Section 1.18.8.3).

1.18.8.1

Stoichiometric b-hydride elimination

Stoichiometric studies of LA-induced b-hydride elimination have primarily focused on the coupling of ethylene and CO2 to produce acrylic acid, a potentially valuable method for utilizing CO2 to produce value-added chemicals.95 The Bernskoetter group has demonstrated that the addition of a LA enables b-hydride elimination from a dppf nickelalactone complex (dppf ¼ 1,10 -bis(diphenylphosphino)ferrocene) by binding to the carbonyl oxygen, weakening the NidO bond, and opening a coordination site for the elimination to occur (Scheme 48).96 After initial b-hydride elimination from a five-membered nickelalactone, the alkene reinserts into the NidH bond to generate a four-membered nickelalactone, with the LA bound to the carbonyl oxygen. Theoretical studies support that the roll of the LA is to facilitate the initial b-hydride elimination from the five-membered nickelalactone.97

Scheme 48 Proposed LA-assisted b-H elimination from a nickelalactone species to form a four-membered metallacycle.

In 2017, Goldman et al. published a study on LA-catalyzed b-hydride elimination from a pincer Ir acetate complex, as well as the microscopic reverse process, alkene insertion.98 The Na+ added was found to aid the elimination reaction by promoting k2-k1 dechelation of the acetate ligand, generating the required vacant coordination site for b-hydride elimination. The accelerating effect was large, with the reaction proceeding to completion in 45 min at −15  C in the presence of NaBArF 4 , while the LA-free reaction was less than 10% complete after 13 h at 125  C (Scheme 49). These studies demonstrate the powerful influence that LA additives can have on the rate of b-hydride elimination.

Scheme 49 Facile b-hydride elimination enabled by a LA.

578

1.18.8.2

Lewis Acid Participation in Organometallic Chemistry

Acrylate synthesis

Limbach et al. described the synthesis of sodium acrylate from CO2, ethylene, and excess base utilizing a Ni catalyst with a bidentate phosphine ligand (Scheme 50).99 Unsuccessful attempts to use an organic base rather than bases with inorganic sodium salts led to the discovery that Na+ is a key component in catalysis. DFT studies determined that the kinetic barrier to b-hydride elimination without the Na+ is 29 kcal/mol, whereas LA binding to the anionic carboxylate group lowers the barrier to 23 kcal/mol. Subsequently, Vogt and co-workers reported a Ni catalyst for acrylate formation that utilized LiCl, NEt3 as the base, and a Zn reductant.100 The proposed mechanism invokes the same coordination of Li+ to stabilize the negative charge buildup on the carboxylate during b-hydride elimination.

Scheme 50 Sodium acrylate synthesis via LA-induced b-hydride elimination.

1.18.8.3

Allene carboxylation

The Iwasawa group reported that the combination of a pincer-supported Pd catalyst and excess LA (AlEt3 or ZnEt2) is capable of facilitating allene carboxylation (Scheme 51A).101 Subsequent mechanistic studies revealed that the LA plays two roles in the reaction: (i) it acts as a transmetallating agent to replace a Pd-carboxylate with a Pd ethyl species;102 and (ii) it assists with b-hydride elimination from a Pd ethyl species to generate a Pd hydride and ethylene (Scheme 51B).

(A)

(B)

Scheme 51 (A) Catalytic allene carboxylation with a pincer-supported Pd complex and (B) the proposed roles of the LA.

1.18.9

1,2-Addition

LAs can facilitate 1,2-addition reactions across metal-ligand bonds. Typically, the electrophilic LA polarizes the substrate to promote activation and enable the addition. LA-assisted 1,2-addition has been observed with a variety of small molecules including H2, CO2, and ethylene. It has also been used as a strategy to promote hydrogenation reactions (Section 1.18.9.2).

1.18.9.1

Stoichiometric 1,2-addition

The Erker group reported H2 addition across a zirconocene bound to a LA-functionalized alkyne (Scheme 52A).103 The reaction is high yielding and occurs under mild conditions. The borane LA acts as a hydride acceptor to facilitate the reaction. Wass and coworkers reported similar cooperative activity between a Pt(0) complex and B(C6F5)3 in the activation of small molecules such as H2, CO2, and ethylene (Scheme 52B).104

Lewis Acid Participation in Organometallic Chemistry

579

(A)

(B)

Scheme 52 (A) Activation of H2 under mild conditions by a zirconocene/LA system. (B) 1,2-Addition of hydrogen using B(C6F5)3 to give a Pt hydride.

1,2-Addition of small molecules such as H2 across EdM bonds, where E is a Lewis acidic Z-type ligand, have been studied extensively.105 The Peters group in particular reported several pioneering examples of reversible H2 activation across NidB, FedB, and CodB bonds in chelating diphosphine-borane complexes.105b A representative example is shown in Scheme 53.106 The mechanism of these reactions has been studied both experimentally and computationally.107 In the proposed pathway, initial coordination of H2 to the metal center generates a dihydrogen complex. The HdH bond is then cleaved in a polarized manner by the electron-deficient borane ligand and electron-rich metal. This process is unique from other more common heterolytic HdH bond cleavage mechanisms involving OM complexes, where the metal is electron-deficient and the ligand electron-rich. In fact, the electrophilicity of LAs that act as Z-type ligands enables facile 1,2-addition reactions of other small molecules such as CO2 and ethylene.

Scheme 53 1,2-Addition of H2 in a Ni diphosphine-borane complex.

1.18.9.2

H2 activation for hydrogenation

Noyori’s catalyst, trans-{(S)-binap}{(S,S)-dpen}RuCl2 (binap ¼ 1,10 -binaphthalene- 2,20 -diylbis(diphenyl-phosphane), dpen ¼ diphenylethylenediamine) requires a Lewis acidic alkali metal for high activity, and is commonly used in the asymmetric hydrogenation of ketones (Scheme 54A).108 Activation by base generates a Ru-amide (Scheme 54B). In conditions with a low concentration of base, this Ru-amide undergoes a 1,2-addition of H2, however the mechanism in the presence of a high concentration of base is more complicated. Specifically, Dub and Gordon propose a mechanism in which the K+ associated with an alkoxide interacts with the amide N atom. This interaction facilitates the facile deprotonation of a molecular H2 complex, which results in the generation of the product alcohol without the formal 1,2-addition of H2. (Scheme 54C).109

(A)

(B)

(C)

Scheme 54 (A) Asymmetric hydrogenation of ketones. (B) Activation of Noyori’s catalyst with KOtBu and 1,2-addition of H2. (C) Mechanism proposed by Dub and Gordon in the presence of a high concentration of base.

580

Lewis Acid Participation in Organometallic Chemistry

The Peters group examined the reversible activation of H2 by Co complexes with pincer ligands containing a central boron atom, which acts as Z-type ligand.107b,110 This scaffold allows for H2 addition across the Co–boryl bond (Scheme 55A), and these complexes are active in alkene hydrogenation (Scheme 55B) and borane dimethylamine dehydrogenation (Scheme 55C).

(A)

(B)

(C)

Scheme 55 (A) Reversible H2 activation across a Co–boryl bond. (b) Catalytic alkene hydrogenation and (C) borane dimethylamine dehydrogenation with a Co pincer complex.

1.18.10

Conclusions

This article demonstrates that LAs can promote a large number of elementary OM reactions, either as exogenous additives or when they are incorporated into the ancillary ligand of an OM complex. In many cases, the use of stoichiometric or catalytic LAs in conjunction with OM catalysts has resulted in more active, stable, and selective catalysts. A number of detailed studies have revealed the underlying reasons why LAs can promote OM reactions, which makes it easier to rationally introduce LAs into a reaction. Nevertheless, a complication in utilizing LAs in OM chemistry is that the reactivity of a specific LA varies as the OM complex is changed. For example, Li+ may enhance OA of a substrate with one metal complex, but will promote the decomposition of another complex. Some reactions require neutral LAs, such as boranes, and others require cationic LAs, such as alkali salts, for improved reactivity to be observed. As a result, it can often be challenging to find a LA that will promote the desired reaction. However, the beneficial effects of LAs in both stoichiometric and catalytic studies suggest that the number of examples of OM reactions utilizing LA additives is likely to increase in future years.

References 1. Lewis, G. N. Valence and the Structure of Atoms and Molecules; Wiley: New York, 1923. 2. Brönsted, J. N. Recl. Trav. Chim. Pays-Bas 1923, 42, 718. 3. (a) Amgoune, A.; Bourissou, D. Chem. Commun. 2011, 47, 859; (b) Braunschweig, H.; Dewhurst, R. D. Dalton Trans. 2011, 40, 549; (c) Parkin, G. Organometallics 2006, 25, 4744. 4. Hill, A. F.; Owen, G. R.; White, A. J. P.; Williams, D. J. Angew. Chem. Int. Ed. 1999, 38, 2759. 5. Bouhadir, G.; Bourissou, D. Chem. Soc. Rev. 2016, 45, 1065. 6. Sivaev, I. B.; Bregadze, V. I. Coord. Chem. Rev. 2014, 270-271, 75. 7. Beckett, M. A.; Strickland, G. C.; Holland, J. R.; Sukumar Varma, K. Polymer 1996, 37, 4629. 8. (a) Huang, J.; Haar, C. M.; Nolan, S. P.; Marcone, J. E.; Moloy, K. G. Organometallics 1999, 18, 297; (b) Nolan, S. P.; Hoff, C. D.; Landrum, J. T. J. Organomet. Chem. 1985, 282, 357. 9. Christe, K. O.; Dixon, D. A.; McLemore, D.; Wilson, W. W.; Sheehy, J. A.; Boatz, J. A. J. Fluor. Chem. 2000, 101, 151. 10. Ghosh, I.; Jacobi, P. A. J. Org. Chem. 2002, 67, 9304. 11. Ohashi, M.; Shibata, M.; Saijo, H.; Kambara, T.; Ogoshi, S. Organometallics 2013, 32, 3631. 12. Fang, X.; Yu, P.; Morandi, B. Science 2016, 351, 832. 13. Fang, X.; Yu, P.; Prina Cerai, G.; Morandi, B. Chem. A Eur. J. 2016, 22, 15629. 14. Ni, S.-F.; Yang, T.-L.; Dang, L. Organometallics 2017, 36, 2746. 15. Bhunia, A.; Bergander, K.; Studer, A. J. Am. Chem. Soc. 2018, 140, 16353. 16. (a) Nakao, Y.; Kanyiva, K. S.; Hiyama, T. J. Am. Chem. Soc. 2008, 130, 2448; (b) Nakao, Y.; Idei, H.; Kanyiva, K. S.; Hiyama, T. J. Am. Chem. Soc. 2009, 131, 15996; (c) Nakao, Y.; Yamada, Y.; Kashihara, N.; Hiyama, T. J. Am. Chem. Soc. 2010, 132, 13666; (d) Nakao, Y. Chem. Rec. 2011, 11, 242; (e) Okumura, S.; Tang, S.; Saito, T.; Semba, K.; Sakaki, S.; Nakao, Y. J. Am. Chem. Soc. 2016, 138, 14699; (f ) Okumura, S.; Nakao, Y. Org. Lett. 2017, 19, 584; (g) Nakao, Y.; Idei, H.; Kanyiva, K. S.; Hiyama, T. J. Am. Chem. Soc. 2009, 131, 5070. 17. Nakao, Y.; Yada, A.; Ebata, S.; Hiyama, T. J. Am. Chem. Soc. 2007, 129, 2428.

Lewis Acid Participation in Organometallic Chemistry

581

18. (a) Nakao, Y.; Ebata, S.; Yada, A.; Hiyama, T.; Ikawa, M.; Ogoshi, S. J. Am. Chem. Soc. 2008, 130, 12874; (b) Nakao, Y.; Hirata, Y.; Tanaka, M.; Hiyama, T. Angew. Chem. Int. Ed. 2008, 47, 385; (c) Hirata, Y.; Tanaka, M.; Yada, A.; Nakao, Y.; Hiyama, T. Tetrahedron 2009, 65, 5037; (d) Hirata, Y.; Yukawa, T.; Kashihara, N.; Nakao, Y.; Hiyama, T. J. Am. Chem. Soc. 2009, 131, 10964; (e) Yada, A.; Yukawa, T.; Nakao, Y.; Hiyama, T. Chem. Commun. 2009, 3931; (f ) Nakao, Y.; Yada, A.; Hiyama, T. J. Am. Chem. Soc. 2010, 132, 10024; (g) Hirata, Y.; Yada, A.; Morita, E.; Nakao, Y.; Hiyama, T.; Ohashi, M.; Ogoshi, S. J. Am. Chem. Soc. 2010, 132, 10070; (h) Koester, D. C.; Kobayashi, M.; Werz, D. B.; Nakao, Y. J. Am. Chem. Soc. 2012, 134, 6544; (i) Minami, Y.; Yoshiyasu, H.; Nakao, Y.; Hiyama, T. Angew. Chem. Int. Ed. 2013, 52, 883; (j) Miyazaki, Y.; Ohta, N.; Semba, K.; Nakao, Y. J. Am. Chem. Soc. 2014, 136, 3732. 19. Frost, G. B.; Serratore, N. A.; Ogilvie, J. M.; Douglas, C. J. J. Org. Chem. 2017, 82, 3721. 20. Pan, Z.; Wang, S.; Brethorst, J. T.; Douglas, C. J. J. Am. Chem. Soc. 2018, 140, 3331. 21. Jayarathne, U.; Zhang, Y.; Hazari, N.; Bernskoetter, W. H. Organometallics 2017, 36, 409. 22. Zhu, B.; Du, G.-F.; Ren, H.; Yan, L.-K.; Guan, W.; Su, Z.-M. Organometallics 2017, 36, 4713. 23. Ren, H.; Du, G.-F.; Zhu, B.; Yang, G.-C.; Yao, L.-S.; Guan, W.; Su, Z.-M. Organometallics 2018, 37, 2594. 24. Cornella, J.; Martin, R. Org. Lett. 2013, 15, 6298. 25. Jia, X.-G.; Guo, P.; Duan, J.; Shu, X.-Z. Chem. Sci. 2018, 9, 640. 26. Nakao, Y.; Morita, E.; Idei, H.; Hiyama, T. J. Am. Chem. Soc. 2011, 133, 3264. 27. Anand, M.; Sunoj, R. B. Org. Lett. 2012, 14, 4584. 28. Tamaki, T.; Ohashi, M.; Ogoshi, S. Angew. Chem. Int. Ed. 2011, 50, 12067. 29. Nakai, K.; Kurahashi, T.; Matsubara, S. Org. Lett. 2013, 15, 856. 30. Guan, W.; Sakaki, S.; Kurahashi, T.; Matsubara, S. ACS Catal. 2015, 5, 1. 31. Lee, J. T.; Thomas, P. J.; Alper, H. J. Org. Chem. 2001, 66, 5424. 32. (a) Getzler, Y. D. Y. L.; Mahadevan, V.; Lobkovsky, E. B.; Coates, G. W. J. Am. Chem. Soc. 2002, 124, 1174; (b) Mahadevan, V.; Getzler, Y. D. Y. L.; Coates, G. W. Angew. Chem. Int. Ed. 2002, 41, 2781; (c) Church, T. L.; Getzler, Y. D. Y. L.; Byrne, C. M.; Coates, G. W. Chem. Commun. 2007, 657. 33. Molnar, F.; Luinstra, G. A.; Allmendinger, M.; Rieger, B. Chem. A Eur. J. 2003, 9, 1273. 34. Church, T. L.; Getzler, Y. D. Y. L.; Coates, G. W. J. Am. Chem. Soc. 2006, 128, 10125. 35. Lei, Y.; Zhang, R.; Han, W.; Mei, H.; Gu, Y.; Xiao, B.; Li, G. Catal. Commun. 2013, 38, 45. 36. Stahl, S. S.; Labinger, J. A.; Bercaw, J. E. J. Am. Chem. Soc. 1995, 117, 9371. 37. Purcell, K. F.; Drago, R. S. J. Am. Chem. Soc. 1966, 88, 919. 38. Shen, Q.; Hartwig, J. F. J. Am. Chem. Soc. 2007, 129, 7734. 39. Driver, M. S.; Hartwig, J. F. J. Am. Chem. Soc. 1997, 119, 8232. 40. (a) Liberman-Martin, A. L.; Bergman, R. G.; Tilley, T. D. J. Am. Chem. Soc. 2013, 135, 9612; (b) Liberman-Martin, A. L.; Levine, D. S.; Liu, W.; Bergman, R. G.; Tilley, T. D. Organometallics 2016, 35, 1064. 41. Liberman-Martin, A. L.; Levine, D. S.; Ziegler, M. S.; Bergman, R. G.; Tilley, T. D. Chem. Commun. 2016, 52, 7039. 42. Taylor, B. W.; Swift, H. E. J. Catal. 1972, 26, 254. 43. (a) Seidel, W. C.; Tolman, C. A. Ann. N. Y. Acad. Sci. 1983, 415, 201; (b) Tolman, C. A.; Seidel, W. C.; Druliner, J. D.; Domaille, P. J. Organometallics 1984, 3, 33; (c) Tolman, C. A.; McKinney, R. J.; Seidel, W. C.; Druliner, J. D.; Stevens, W. R. In Advances in Catalysis; Eley, D. D., Pines, H., Weisz, P. B., Eds.; Academic Press, 1985; vol. 33; p 1. 44. Bini, L.; Pidko, E. A.; Müller, C.; van Santen, R. A.; Vogt, D. Chem. A Eur. J. 2009, 15, 8768. 45. Shiba, T.; Kurahashi, T.; Matsubara, S. J. Am. Chem. Soc. 2013, 135, 13636. 46. Duez, S.; Steib, A. K.; Manolikakes, S. M.; Knochel, P. Angew. Chem. Int. Ed. 2011, 50, 7686. 47. West, N. M.; Miller, A. J. M.; Labinger, J. A.; Bercaw, J. E. Coord. Chem. Rev. 2011, 255, 881. 48. Collman, J. P.; Cawse, J. N.; Brauman, J. I. J. Am. Chem. Soc. 1972, 94, 5905. 49. Collman, J. P.; Finke, R. G.; Cawse, J. N.; Brauman, J. I. J. Am. Chem. Soc. 1978, 100, 4766. 50. (a) Butts, S. B.; Holt, E. M.; Strauss, S. H.; Alcock, N. W.; Stimson, R. E.; Shriver, D. F. J. Am. Chem. Soc. 1979, 101, 5864; (b) Butts, S. B.; Strauss, S. H.; Holt, E. M.; Stimson, R. E.; Alcock, N. W.; Shriver, D. F. J. Am. Chem. Soc. 1980, 102, 5093; (c) Stimson, R. E.; Shriver, D. F. Inorg. Chem. 1980, 19, 1141; (d) Richmond, T. G.; Basolo, F.; Shriver, D. F. Inorg. Chem. 1982, 21, 1272; (e) Richmond, T. G.; Basolo, F.; Shriver, D. F. Organometallics 1982, 1, 1624. 51. (a) Labinger, J. A.; Miller, J. S. J. Am. Chem. Soc. 1982, 104, 6856; (b) Grimmett, D. L.; Labinger, J. A.; Bonfiglio, J. N.; Masuo, S. T.; Shearin, E.; Miller, J. S. J. Am. Chem. Soc. 1982, 104, 6858; (c) Labinger, J. A.; Bonfiglio, J. N.; Grimmett, D. L.; Masuo, S. T.; Shearin, E.; Miller, J. S. Organometallics 1983, 2, 733; (d) Grimmett, D. L.; Labinger, J. A.; Bonfiglio, J. N.; Masuo, S. T.; Shearin, E.; Miller, J. S. Organometallics 1983, 2, 1325. 52. (a) Miller, A. J. M.; Labinger, J. A.; Bercaw, J. E. J. Am. Chem. Soc. 2008, 130, 11874; (b) Miller, A. J. M.; Labinger, J. A.; Bercaw, J. E. Organometallics 2010, 29, 4499; (c) Miller, A. J. M.; Labinger, J. A.; Bercaw, J. E. J. Am. Chem. Soc. 2010, 132, 3301; (d) West, N. M.; Labinger, J. A.; Bercaw, J. E. Organometallics 2011, 30, 2690; (e) Hazari, A.; Labinger, J. A.; Bercaw, J. E. Angew. Chem. Int. Ed. 2012, 51, 8268. 53. (a) Wang, H.-K.; Choi, H. W.; Muetterties, E. L. Inorg. Chem. 1981, 20, 2661; (b) Choi, H. W.; Muetterties, E. L. Inorg. Chem. 1981, 20, 2664. 54. Hazari, N.; Heimann, J. E. Inorg. Chem. 2017, 56, 13655. 55. (a) Darensbourg, D. J.; Pala, M. J. Am. Chem. Soc. 1985, 107, 5687; (b) Miller, A. J. M.; Labinger, J. A.; Bercaw, J. E. Organometallics 2011, 30, 4308; (c) Heimann, J. E.; Bernskoetter, W. H.; Hazari, N.; Mayer, J. M. Chem. Sci. 2018, 9, 6629; (d) Heimann, J. E.; Bernskoetter, W. H.; Hazari, N. J. Am. Chem. Soc. 2019, 141, 10520. 56. (a) Savéant, J.-M. Chem. Rev. 2008, 108, 2348; (b) Chakraborty, S.; Bhattacharya, P.; Dai, H.; Guan, H. Acc. Chem. Res. 2015, 48, 1995; (c) Bernskoetter, W. H.; Hazari, N. Acc. Chem. Res. 2017, 50, 1049. 57. Bontemps, S. Coord. Chem. Rev. 2016, 308, 117. 58. (a) Chakraborty, S.; Zhang, J.; Krause, J. A.; Guan, H. J. Am. Chem. Soc. 2010, 132, 8872; (b) Bontemps, S.; Vendier, L.; Sabo-Etienne, S. Angew. Chem. Int. Ed. 2012, 51, 1671; (c) Espinosa, M. R.; Charboneau, D. J.; Garcia de Oliveira, A.; Hazari, N. ACS Catal. 2019, 9, 301. 59. Chen, J.; McGraw, M.; Chen, E. Y.-X. ChemSusChem 2019, 12, 4543. 60. Mitton, S. J.; Turculet, L. Chem. A Eur. J. 2012, 18, 15258. 61. Jiang, Y.; Blacque, O.; Fox, T.; Berke, H. J. Am. Chem. Soc. 2013, 135, 7751. 62. Bielinski, E. A.; Lagaditis, P. O.; Zhang, Y.; Mercado, B. Q.; Wurtele, C.; Bernskoetter, W. H.; Hazari, N.; Schneider, S. J. Am. Chem. Soc. 2014, 136, 10234. 63. Bielinski, E. A.; Förster, M.; Zhang, Y.; Bernskoetter, W. H.; Hazari, N.; Holthausen, M. C. ACS Catal. 2015, 5, 2404. 64. Nielsen, M.; Alberico, E.; Baumann, W.; Drexler, H.-J.; Junge, H.; Gladiali, S.; Beller, M. Nature 2013, 495, 85. 65. Wei, Z.; de Aguirre, A.; Junge, K.; Beller, M.; Jiao, H. Cat. Sci. Technol. 2018, 8, 3649. 66. Li, Y.; Topf, C.; Cui, X.; Junge, K.; Beller, M. Angew. Chem. Int. Ed. 2015, 54, 5196. 67. Cui, X.; Li, Y.; Topf, C.; Junge, K.; Beller, M. Angew. Chem. Int. Ed. 2015, 54, 10596. 68. Cabrero-Antonino, J. R.; Alberico, E.; Junge, K.; Junge, H.; Beller, M. Chem. Sci. 2016, 7, 3432. 69. Yuan, M.-L.; Xie, J.-H.; Zhou, Q.-L. ChemCatChem 2016, 8, 3036. 70. Zou, Y.-Q.; Chakraborty, S.; Nerush, A.; Oren, D.; Diskin-Posner, Y.; Ben-David, Y.; Milstein, D. ACS Catal. 2018, 8, 8014. 71. Wick, D. D.; Goldberg, K. I. J. Am. Chem. Soc. 1997, 119, 10235. 72. Beswick, C. L.; Marks, T. J. J. Am. Chem. Soc. 2000, 122, 10358. 73. Boudreau, J.; Fontaine, F.-G. Organometallics 2011, 30, 511.

582 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. 105. 106. 107. 108. 109. 110.

Lewis Acid Participation in Organometallic Chemistry Yang, X.; Stern, C. L.; Marks, T. J. J. Am. Chem. Soc. 1994, 116, 10015. Deck, P. A.; Beswick, C. L.; Marks, T. J. J. Am. Chem. Soc. 1998, 120, 1772. Chen, Y.-X.; Metz, M. V.; Li, L.; Stern, C. L.; Marks, T. J. J. Am. Chem. Soc. 1998, 120, 6287. Sattler, A.; Labinger, J. A.; Bercaw, J. E. Organometallics 2013, 32, 6899. Sattler, A.; VanderVelde, D. G.; Labinger, J. A.; Bercaw, J. E. J. Am. Chem. Soc. 2014, 136, 10790. (a) Richmond, T. G.; Shriver, D. F. Organometallics 1983, 2, 1061; (b) Richmond, T. G.; Shriver, D. F. Organometallics 1984, 3, 305. (a) Harrison, D. J.; Lee, G. M.; Leclerc, M. C.; Korobkov, I.; Baker, R. T. J. Am. Chem. Soc. 2013, 135, 18296; (b) Giffin, K. A.; Harrison, D. J.; Korobkov, I.; Baker, R. T. Organometallics 2013, 32, 7424; (c) Andrella, N. O.; Sicard, A. J.; Gorelsky, S. I.; Korobkov, I.; Baker, R. T. Chem. Sci. 2015, 6, 6392. Becica, J.; Dobereiner, G. E. ACS Catal. 2017, 7, 5862. Brereton, K. R.; Smith, N. E.; Hazari, N.; Miller, A. J. M. Chem. Soc. Rev. 2020, 49, 7929. Wiedner, E. S.; Chambers, M. B.; Pitman, C. L.; Bullock, R. M.; Miller, A. J. M.; Appel, A. M. Chem. Rev. 2016, 116, 8655. Ilic, S.; Alherz, A.; Musgrave, C. B.; Glusac, K. D. Chem. Soc. Rev. 2018, 47, 2809. Sarker, N.; Bruno, J. W. J. Am. Chem. Soc. 1999, 121, 2174. (a) Kita, M. R.; Miller, A. J. M. J. Am. Chem. Soc. 2014, 136, 14519; (b) Grajeda, J.; Kita, M. R.; Gregor, L. C.; White, P. S.; Miller, A. J. M. Organometallics 2016, 35, 306; (c) Smith, J. B.; Kerr, S. H.; White, P. S.; Miller, A. J. M. Organometallics 2017, 36, 3094; (d) Camp, A. M.; Kita, M. R.; Grajeda, J.; White, P. S.; Dickie, D. A.; Miller, A. J. M. Inorg. Chem. 2017, 56, 11141. Kita, M. R.; Miller, A. J. M. Angew. Chem. Int. Ed. 2017, 56, 5498. (a) Gregor, L. C.; Grajeda, J.; Kita, M. R.; White, P. S.; Vetter, A. J.; Miller, A. J. M. Organometallics 2016, 35, 3074; (b) Gregor, L. C.; Grajeda, J.; White, P. S.; Vetter, A. J.; Miller, A. J. M. Cat. Sci. Technol. 2018, 8, 3133. (a) Jiang, Y.; Schirmer, B.; Blacque, O.; Fox, T.; Grimme, S.; Berke, H. J. Am. Chem. Soc. 2013, 135, 4088; (b) Jiang, Y.; Huang, W.; Schmalle, H. W.; Blacque, O.; Fox, T.; Berke, H. Organometallics 2013, 32, 7043; (c) Jiang, Y.; Huang, W.; Schmalle, H. W.; Blacque, O.; Fox, T.; Berke, H. Eur. J. Inorg. Chem. 2014, 2014, 140. Ouyang, G.-H.; He, Y.-M.; Li, Y.; Xiang, J.-F.; Fan, Q.-H. Angew. Chem. Int. Ed. 2015, 54, 4334. Drake, J. L.; Manna, C. M.; Byers, J. A. Organometallics 2013, 32, 6891. Zhang, Y.; MacIntosh, A. D.; Wong, J. L.; Bielinski, E. A.; Williard, P. G.; Mercado, B. Q.; Hazari, N.; Bernskoetter, W. H. Chem. Sci. 2015, 6, 4291. (a) Spentzos, A. Z.; Barnes, C. L.; Bernskoetter, W. H. Inorg. Chem. 2016, 55, 8225; (b) Mills, M. R.; Barnes, C. L.; Bernskoetter, W. H. Inorg. Chem. 2018, 57, 1590. Schneidewind, J.; Adam, R.; Baumann, W.; Jackstell, R.; Beller, M. Angew. Chem. Int. Ed. 2017, 56, 1890. Hollering, M.; Dutta, B.; Kühn, F. E. Coord. Chem. Rev. 2016, 309, 51. (a) Jin, D.; Schmeier, T. J.; Williard, P. G.; Hazari, N.; Bernskoetter, W. H. Organometallics 2013, 32, 2152; (b) Jin, D.; Williard, P. G.; Hazari, N.; Bernskoetter, W. H. Chem. A Eur. J. 2014, 20, 3205. (a) Plessow, P. N.; Weigel, L.; Lindner, R.; Schäfer, A.; Rominger, F.; Limbach, M.; Hofmann, P. Organometallics 2013, 32, 3327; (b) Plessow, P. N.; Schäfer, A.; Limbach, M.; Hofmann, P. Organometallics 2014, 33, 3657; (c) Yang, G.; Schäffner, B.; Blug, M.; Hensen, E. J. M.; Pidko, E. A. ChemCatChem 2014, 6, 800. Gao, Y.; Guan, C.; Zhou, M.; Kumar, A.; Emge, T. J.; Wright, A. M.; Goldberg, K. I.; Krogh-Jespersen, K.; Goldman, A. S. J. Am. Chem. Soc. 2017, 139, 6338. Lejkowski, M. L.; Lindner, R.; Kageyama, T.; Bódizs, G.É.; Plessow, P. N.; Müller, I. B.; Schäfer, A.; Rominger, F.; Hofmann, P.; Futter, C.; Schunk, S. A.; Limbach, M. Chem. A Eur. J. 2012, 18, 14017. Hendriksen, C.; Pidko, E. A.; Yang, G.; Schäffner, B.; Vogt, D. Chem. A Eur. J. 2014, 20, 12037. Takaya, J.; Iwasawa, N. J. Am. Chem. Soc. 2008, 130, 15254. Suh, H.-W.; Guard, L. M.; Hazari, N. Chem. Sci. 2014, 5, 3859. Podiyanachari, S. K.; Fröhlich, R.; Daniliuc, C. G.; Petersen, J. L.; Mück-Lichtenfeld, C.; Kehr, G.; Erker, G. Angew. Chem. Int. Ed. 2012, 51, 8830. Forrest, S. J. K.; Clifton, J.; Fey, N.; Pringle, P. G.; Sparkes, H. A.; Wass, D. F. Angew. Chem. Int. Ed. 2015, 54, 2223. (a) Owen, G. R. Chem. Soc. Rev. 2012, 41, 3535; (b) Owen, G. R. Chem. Commun. 2016, 52, 10712. Harman, W. H.; Peters, J. C. J. Am. Chem. Soc. 2012, 134, 5080. (a) Zeng, G.; Sakaki, S. Inorg. Chem. 2013, 52, 2844; (b) Harman, W. H.; Lin, T.-P.; Peters, J. C. Angew. Chem. Int. Ed. 2014, 53, 1081; (c) Ganguly, G.; Malakar, T.; Paul, A. ACS Catal. 2015, 5, 2754. (a) Doucet, H.; Ohkuma, T.; Murata, K.; Yokozawa, T.; Kozawa, M.; Katayama, E.; England, A. F.; Ikariya, T.; Noyori, R. Angew. Chem. Int. Ed. 1998, 37, 1703; (b) Noyori, R.; Ohkuma, T. Angew. Chem. Int. Ed. 2001, 40, 40. Dub, P. A.; Henson, N. J.; Martin, R. L.; Gordon, J. C. J. Am. Chem. Soc. 2014, 136, 3505. Lin, T.-P.; Peters, J. C. J. Am. Chem. Soc. 2013, 135, 15310.

1.19

Organometallic Chemistry on Oxide Surfaces

Matthew P Conley, Jiaxin Gao, Winn Huynh, Jessica Rodriguez, and Kavyasripriya K Samudrala, Department of Chemistry, University of California, Riverside, CA, United States © 2022 Elsevier Ltd. All rights reserved.

1.19.1 Introduction 1.19.2 Oxide supports 1.19.2.1 Surface chemistry of silica 1.19.2.2 Surface chemistry of alumina 1.19.2.3 Surface chemistry of sulfated oxides 1.19.3 Reactions of organometallics with oxides by protonolysis 1.19.3.1 Reactions of organometallics with partially dehydroxylated silica 1.19.3.2 Reactions of organometallics with dOH groups present on sulfated oxide surfaces 1.19.3.3 Factors affecting formation of ^SidOdM or [M][oxide] 1.19.4 Reactivity of strained ^SidOdSi^ bridges on dehydroxylated silica surfaces 1.19.5 Alkyl abstraction by Lewis sites on Al2O3 1.19.6 Heterolytic activation of CdH bonds on Al2O3 1.19.7 Conclusion Acknowledgment References

1.19.1

583 584 585 586 588 590 590 595 597 599 601 603 604 604 604

Introduction

Commodity chemicals (fuels, plastics, NH3, H2SO4, etc.) are manufactured on enormous scales, are major consumers of energy, and contribute directly or indirectly to 35% of GDP.1 These chemicals require an appropriate catalyst to mediate the formation of the desired product. Indeed, 90% of commodity and specialty chemical synthesis involves at least one catalytic step, nearly all of them using a heterogeneous catalyst.2 By necessity these catalysts must be synthetically accessible on large scale and usually involve a combination of co-precipitation, impregnation, vapor deposition, and/or thermal annealing.3–7 Industrial olefin metathesis catalysts are useful examples that show the challenges in accessing heterogeneous catalysts that have similar rate and selectivity profiles as homogeneous catalysts for this reaction. The ethenolysis of 2-butene to form propene using the WO3/SiO2 catalyst is the largest scale application of olefin metathesis, developed in the 1960s by Phillips Petroleum.8,9 WO3/ SiO2 is prepared by heating silica and ammonium metatungstate under air and contains a complex mixture of tungsten sites, some of which are active in propene metathesis above 400  C.10 The temperatures required for metathesis using WO3/SiO2 are significantly higher than temperatures needed to mediate olefin metathesis reactions with homogeneous Mo, W, or Ru catalysts,11,12 and obviously incompatible with functional groups. The active site in WO3/SiO2 is unknown, but is probably a tungsten-oxo-alkylidene bound to the silica surface, shown in Fig. 1A with the initial rate for propene metathesis at 400  C. In contrast, (HMTO)2W(] CHtBu)(]O) (HMTO ¼ 2,6-dimesitylphenoxide, Fig. 1B), a homogeneous mimic of the putative active site in WO3/SiO2, is several orders of magnitude more active than WO3/SiO2, catalyze metathesis at room temperature, and are compatible with functionalized olefins.13 These results pose a simple question, are alkylidenes supported on silica inherently unreactive in olefin metathesis reactions? This is easy to test by reacting an alkylidene complex with partially dehydroxylated silica, and in fact many well-defined alkylidenes supported on silica are more active in olefin metathesis than the homogeneous analogues.14–17 For example, grafting (HMTO)2W(]CHtBu)(]O) (Fig. 1B) onto silica generates (^SiO)W(]O)(]CHtBu)(OTMH) (Fig. 1C), a well-defined tungsten-oxo-alkylidene analogue of the putative active site in WO3/SiO2.18,19 This material is at least 100 times more active than (HMTO)2W(]CHtBu)(]O) at room temperature in metathesis of cis-4-nonene, and several orders of magnitude more active than WO3/SiO2. Related examples containing organometallic tungsten supported on silica are also active in propene metathesis at much lower temperatures than WO3/SiO2.20–23 This discussion shows that incorporation of well-defined organometallics onto oxides is a viable strategy to achieve the goal of more efficient and selective heterogeneous catalysts. Formation of well-defined organometallics on heterogeneous supports is surface organometallic chemistry (SOMC).24–32 This methodology provides opportunities to rationally tune the structure of catalytic sites on surfaces to affect catalytic properties, the most common methodology to optimize homogeneous catalysts, and a long-term goal in heterogeneous catalysis. The number of book chapters,33–35 reviews,36–42 perspectives,43–49 and accounts50,51 published in this field since 2015 is a testimony to the rapid growth of SOMC. This is due in large part to the availability of advanced characterization techniques that are accelerating the development of new well-defined heterogeneous catalysts, the most important being solid-state NMR spectroscopy,52 dynamic nuclear polarization (DNP) enhanced NMR spectroscopy,53,54 and X-ray adsorption spectroscopy (XAS).55 Simulation of these spectroscopic observables using DFT calculations of organometallics on surfaces provides additional support for the well-defined coordination environment in these materials.56

Comprehensive Organometallic Chemistry IV

https://doi.org/10.1016/B978-0-12-820206-7.00037-8

583

584

Organometallic Chemistry on Oxide Surfaces

(A)

(B)

(C)

Fig. 1 Rate of ethenolysis of 2-butene at 400  C for WO3/SiO2 (A). The alkylidene structure shown for WO3/SiO2 in (A) has not been observed, but is suggested based on the extensive homogeneous chemistry of W(VI) alkylidenes. Structure of (HMTO)2W(]CHtBu)(]O) (HMTO ¼ 2,6-dimesitylphenoxide) (B), and (^SiO)W(]O)(]CHtBu)(OTMH) (C). The TOFini shown in (B) and (C) are for metathesis of cis-4-nonene.

Fig. 2 Reaction pathways to form well defined organometallics on oxides.

Previous reviews in SOMC focus on the structure and function of the well-defined organometallic on an oxide. The aim of this chapter is to describe how oxide surfaces affect the structure of supported organometallics, essentially a description of surface ligand effects. Organometallics react with oxides by one of the four reactions shown in Fig. 2. The most common method is to react a partially dehydroxylated oxide, nearly all which are terminated with dOH groups, with a generic LnMdCH3 to form either ^EOdMLn or [MLn][OE^] ion-pairs. Reactions of LnMdCH3 with strong Lewis sites on oxides to form [LnM][Rdoxide] are analogous to reactions of Lewis acids (e.g. B(C6F5)3, [Ph3C][B(C6F5)4]) to form ion pairs in solution.57 This pathway is less common that protonolysis, but is important with alumina supports. Less common grafting pathways involve the electrophilic opening of strained ^EdOdE^ bridges by LnMdCH3 to form ^EdOMLn and H3CdE^, or heterolytic cleavage of a CdH bond in LnMdCH3 across an ^EdOdE^ bridge to form ^ECH2MLn species.

1.19.2

Oxide supports

The most common method to optimize catalytic properties of a homogeneous catalyst is to systematically modify the sterics and electronics around the active site, typically a metal. There are rigorous models to quantify these properties (nCO/cone angle,58 buried volume,59,60 bite angle,58 etc.) in transition metal complexes. Much like a chemist would design a ligand to affect rate or selectivity based on structure property trends, or on chemical intuition, those interested in SOMC must first consider the nature of the inorganic support. The greatest challenge in SOMC is understanding how an inorganic support, usually a partially dehydroxylated oxide, interacts with an organometallic (Fig. 2). Below are descriptions of silica, alumina, and sulfated oxides. These are common oxides used in SOMC, and serve as reference surfaces for the organometallic chemistry described in the following sections. The SOMC of less common oxide supports was recently reviewed.49

Organometallic Chemistry on Oxide Surfaces

1.19.2.1

585

Surface chemistry of silica

High surface area amorphous silica is thermally and mechanically stable, a mild Brønsted acid, and does not engage in significant redox chemistry. These features explain why amorphous silica is a very common support in heterogeneous catalysis, and by far the most common support in SOMC. Amorphous silica is composed of connected SiO4 tetrahedra that form the bulk SiO2 network and is terminated by SidOH groups. The bulk SiO2 network is complex. For example, vitreous thin SiO2 films grown on Ru(0001) surfaces shown in Fig. 3 are composed of connected SiO4 tetrahedra forming between 4-membered and 9-membered rings, with (SiO)3 being the most dominant ring structure (Fig. 3A–C).61 Detailed analysis of bond angles in the ring structures show that the most common SidOdSi bond angles are the predicted tetrahedral values (109.5 , Fig. 3D) and that the most common SidOdSi bond angles in this sample are close to those observed experimentally in vitreous silica (Fig. 3E).62,63 Though the bulk SiO2 network is undeniably complex, simple models describing the surface chemistry of silica have predictive power.64 Under ambient conditions amorphous silica contains several monolayers of physisorbed water that can be removed by heating to 200  C under vacuum. This treatment exposes ^SidOdSi^ siloxane bridges and the three types of surface silanols shown in Fig. 4A. Partial dehydroxylation of silica occurs at higher temperatures, which results in condensation of nearby silanols to generate water and new ^SidOdSi^ siloxane bridges, shown for vicinal silanols in Fig. 4B. Fig. 4C shows the surface coverage (Y) of silanols on silica treated at different dehydroxylation temperatures. There are several methods to determine Y for oxides, though the most common to practitioners of SOMC is to quantify the amount of alkane generated in the reaction of an oxide with a Grignard or alkyllithium reagent in hydrocarbon slurry. As the dehydroxylation temperature increases Y decreases. Silica partially dehydroxylated at 700  C (SiO2-700) is the most common support in SOMC and contains 1 dOH site nm−2. These trends are also reflected in the FTIR spectrum of partially dehydroxylated silica, shown in Fig. 4D for SiO2-300 (top) and SiO2-700 (bottom). SiO2-300 contains  2 dOH nm−2, and the FTIR spectrum contains a sharp nOH at 3747 cm−1 assigned to isolated ^SidOH and geminal ]Si(OH)2, and a broad signal centered at 3650 cm−1 for hydrogen bonded vicinal silanols. The FTIR spectrum of SiO2-700 lacks the broad feature and contains only a sharp nOH at 3747 cm−1, corresponding to 90% isolated ^SidOH and 10% geminal ]Si(OH)2.64 This model for partial dehydroxylation of silica is also supported by solid-state NMR studies. 1H Magic Angle Spinning (MAS) NMR spectra contain narrow ^SidOH signals at high dehydroxylation temperatures, and 29Si MAS NMR spectra contain signals assigned to ^SidOH (d 29Si ¼ −99 ppm) and ]Si(OH)2 (d 29Si ¼ − 89 ppm), which decrease when silica is heated to progressively higher temperatures65–67; consistent with a depletion of silanols on the silica surface following dehydroxylation.

Fig. 3 STM image of vitreous silica on Ru(0001) (A) and a model showing Si atoms (green) and oxygen atoms (red, B). Histogram of the different ring sizes from the STM image (C). SidO bond angles in SiO4 tetrahedra (D) and SidOdSi bond angles connecting SiO4 tetrahedra (E) from STM data. Adapted with permission from reference Lichtenstein, L.; Büchner, C.; Yang, B.; Shaikhutdinov, S.; Heyde, M.; Sierka, M.; Włodarczyk, R.; Sauer, J.; Freund, H.-J. Angew. Chem., Int. Ed. 2012, 51, 404–407.

586

Organometallic Chemistry on Oxide Surfaces

(A)

(B)

(C)

Fig. 4 Silanols present on the silica surface (A), condensation of vicinal silanols to release water and form a new ^SidOdSi^ (B), Y of silanols on silica treated at different temperatures (C), FTIR of SiO2-300 and SiO2-700 (D).

More complex 17O NMR experiments also support this model. 17O is NMR active and low natural abundance (0.037%), though enrichment is fairly straightforward and economical using vapor phase enrichment techniques.68,69 The 17O NMR spectrum of enriched silica contains separated signals for Sid OH and Sid OdSi. 1Hd17O J-HMQC experiments select 17O signals that have scalar (J) coupling to 1H nuclei. The 1Hd17O J-HMQC spectra of SiO2-700 contains a narrow correlation between the 17O signal for Sid OH and Sid OH, Fig. 5A, indicating that the isolated ^SidOH are in well-defined environments and that these silanols are not J-coupled to ^SidOdSi^ bridges.70 However, the spectrum of SiO∗ has a much broader 2 -200 1 Hd17O correlation due to the network of H-bonded silanols present in this material. This experiment shows the sensitivity of 17 O NMR to local structure, and also shows the differences in the silanol structure for silica surfaces treated at different dehydroxylation temperatures. DFT is indispensable in understanding the structures of well-defined sites supported on oxides. Simple small cluster models (e.g. H3SidOH, (MeO)3SidOH, polysilsesquioxanes) that contain an isolated SidOH group are dramatic oversimplifications for the silica surface, but are useful to predict reactivity or spectroscopic trends of supported species on silica.71–77 Amorphous silica obviously does not contain a repeating unit cell, which until recently limited periodic DFT studies of organometallics supported on silica. Early calculations on periodic silica surfaces used the crystalline silica cristobalite as a model for isolated silanols.78 More realistic periodic models of amorphous silica that closely match experimental FTIR properties containing 1.1–7.2 dOH nm−2 are now available.79,80

1.19.2.2

Surface chemistry of alumina

Aluminum oxide is more complicated than silica. Several stable phases (a, d, g, Z and y) of alumina are available81,82 with g-alumina as the most common phase used in SOMC. Depending on the supplier or synthetic method, g-alumina has moderate surface area (200 m2 g−1), is thermally stable to 700  C under vacuum for 4–12 h,83 but can undergo phase transition to d/y-alumina at higher temperatures, a process facilitated by steam.84 The aluminum sites present of the g-alumina surface range from tri-coordinate (AlIII) to hexa-coordinate (AlVI), and the surface is terminated with dOH groups. The nature of the dOH groups on g-alumina are different than the silanols present on SiO2. Fig. 6A shows the nOH region of the FTIR spectra of g-alumina at various dehydroxylation temperatures. The nOH stretches are assigned to Al −OH groups bound to various aluminum sites, or m2- or m3-OH bridging Al sites shown in Fig. 6B, and this model is supported by solid-state NMR studies.85 The lineshape of the nOH stretches on g-alumina do not change appreciably at higher dehydroxylation temperatures, in contrast to what is observed with SiO2. This result indicates that all of the AldOH sites in Fig. 6B are present at all dehydroxylation temperatures. The dehydroxylation process is also accompanied by the formation of Lewis acidic-Al sites in g-alumina treated above 400  C that bind pyridine, CO, N2, and heterolytically cleave XdH bonds (X ¼ H, CH3).83,86–89

Organometallic Chemistry on Oxide Surfaces

587

Fig. 5 1Hd17O J-HMQC MAS NMR spectra of SiO∗ (A); 1Hd17O J-HMQC MAS NMR spectra of SiO∗ (B). The arrows connect the types of silanols present 2 -700 2 -200 ∗ ∗ and SiO-700 shown above the spectra. Reproduced with permission from reference Merle, N.; Trébosc, J.; Baudouin, A.; Rosal, I. D.; Maron, L.; in SiO-200 2 2 Szeto, K.; Genelot, M.; Mortreux, A.; Taoufik, M.; Delevoye, L.; Gauvin, R. M. J. Am. Chem. Soc. 2012, 134, 9263–9275. (A)

1

(B) 3673

0.8 3629

3725

abs.

0.6

3771

0.4

(C) 350°C

3795 400°C

0.2 700°C

500°C

850°C 1000°C

0 3800

3700

3600

3500 –1

Wavenumber / cm

Fig. 6 FTIR spectra of g-Al2O3 after different temperature treatments (A). dOH groups present on the g-Al2O3 surface with FTIR assignment (B). Y of dOH groups (in OH nm−2) and surface area (in m2 g−1) of g-Al2O3 as a function of temperature (C). Reproduced with permission from reference Copéret, C.; Comas-Vives, A.; Conley, M. P.; Estes, D. P.; Fedorov, A.; Mougel, V.; Nagae, H.; Núñez-Zarur, F.; Zhizhko, P. A. Chem. Rev. 2016, 116, 323–421.

588

Organometallic Chemistry on Oxide Surfaces

(A)

(B)

AllVa'

(C)

1.88

O2a' Allll

Allll O2a

O3a'

AllVa O2b

O3a O3b

AllVb

1.60 O2a'

AllVa 1.88 AllVb

O3a'

AllVa Allll

O2a O2b

O3a O3b

AllVb

Fig. 7 Fully dehydroxylated 110 surface of g-Al2O3 (A). 110 surface of g-Al2O3 containing 3 OH nm−2, water interacting with AlIII sites (B) or AlIV sites (C). The dashed line is the unit cell for these calculations. Reproduced with permission from reference Wischert, R.; Laurent, P.; Coperet, C.; Delbecq, F.; Sautet, P. J. Am. Chem. Soc. 2012, 134, 14430–14449..

Similar to SiO2, increasing dehydroxylation temperatures result in decreasing Y for AldOH sites (Fig. 6C). For example, Al2O3contains 0.7 OH nm−2. Fully dehydroxylated alumina is accessible at 1000  C, but d and y phase contamination is unavoidable at these temperatures from powder XRD diffraction.86 DFT models of g-alumina surfaces are based on the dehydration of boehmite containing 25% tetrahedral and 75% octahedral aluminum sites.88 The completely dehydrated (110) surface, which is the experimentally more accessible termination in g-aluminas, contains one tricoordinate trigonal planar AlIII and two tetracoordinate AlIV sites, Fig. 7 (left). The fully dehydrated (110) surface model is metastable, and from the properties of g-alumina at very high temperatures required for a dehydrated surface, not experimentally realistic. Adsorption of one molecule of water to the fully dehydrated unit cell ( 3 dOH nm−2) results in exergonic heterolytic splitting of an HdO bond across the AlIIIdO bond to form an AldOH and Ald(OH)dAl site (−53.1 kcal mol−1), Fig. 7 (middle). Heterolytic cleavage of water across two tetrahedral aluminum sites, shown in Fig. 7 (right), is only 11 kcal mol−1 less stable, suggesting that surfaces containing free AlIII sites are accessible. The AlIII “defect” site is proposed to result in the reactivity of alumina with probe molecules and is also suggested to promote carbon-carbon bond formation in the presence of CH3F and Me2O with the participation of the Lewis acid sites.89,90 The surface chemistry of silica and g-alumina is obviously quite different. Unlike silica, the g-alumina surface does not contain a network of AldOH sites that undergo condensation reactions, but rather a network of AldOdAl sites that react with water by heterolytic cleavage to form AldOH sites. The understanding of this process from the DFT models implies that the Lewis sites on alumina play a role in reactions of g-alumina with organometallics, which will be described below, and underlies the importance of understanding basic chemistry of the surface when designing reactions in SOMC. Though less common in SOMC, silica coated alumina, or crystalline silica-alumina materials (e.g. zeolites), are extremely important in the petrochemical industry. In general, silica aluminas are prepared by co-precipitation,91 sol-gel,92–94 thermolysis of molecular precursors,95–97 or vapor deposition.98,99 Acetonitrile, pyridine, and CO adsorb to silica-alumina, indicating that Lewis acidic Al-sites are present.100–104 The silica-alumina surface contains isolated silanols,97 but the nOH FTIR signature is highly dependent on the synthetic methods used to make the material, and in some cases OH groups bonded to Al centers are observed in IR spectroscopy.104,105 The bridging silanol shown in Fig. 8 was characterized in detail by DNP NMR spectroscopy and studied by periodic DFT methods,106 and is source of Bronsted acidity in silica alumina. 700

1.19.2.3

Surface chemistry of sulfated oxides

Sulfated metal oxides (SMOs) are another common support traditionally encountered in applications for olefin polymerization or arene hydrogenation.25,26 These materials are prepared by contacting a native oxide with dilute sulfuric acid, followed by high temperature calcination. The SMO surface is generally assumed to contain dOH groups that are significantly more acidic than isolated silanols on silica. Initial reports of sulfated zirconium oxide (SZO) showed that Brønsted sites on this material isomerize n-butane at lower temperatures than neat H2SO4, suggesting that this material is superacidic.107 SZO was suggested to protonate very weak colorimetric Hammett bases (H0 < −16), consistent with this interpretation.108 However, colorimetric protonation studies using Hammett indicators on oxides can be misleading.109 Subsequent studies suggested that butane isomerization occurs due to trace butene in the reaction feed,110 and that pyrosulfate (e.g. oxidative) sites are responsible for isomerization activity in purified feeds lacking butene.111

Fig. 8 The bridging silanol in silica alumina.

Organometallic Chemistry on Oxide Surfaces

589

(A)

(B)

Fig. 9 Reaction of phosphines with SZO to form [R3PH][SZO] sites (A). This Langmuir binding equation is given below the reaction. Thermodynamic cycle explaining the acidity of the dOH site in SZO (B).

Experimental studies are inconsistent with the presence of superacidic dOH groups on SZO surfaces. SZO reacts with trialkylphosphines to form [R3PH][SZO] ion pairs (Fig. 9).112 (tBu)2ArP, where the para position of the Ar group contains substituents that donate or withdraw electron density spanning  4 pKa units, were chosen to evaluate how the electronics at phosphorous affects formation of [(tBu)2ArPH][SZO]. In all cases, MeCN slurries of tBu2PAr bind to the acidic dOH sites on SZO and follow simple Langmuir isotherms, which relate surface coverage to Keq. Evaluation of Keq for each phosphine showed that phosphines containing electron withdrawing substituents bind weaker to SZO than more electron rich phosphines, as expected for weak acids interacting with weak bases. In addition, the dOH sites on SZO follow linear free energy relationships across this series, as expected for Hammett acids. Finally, the dOH sites on SZO are unable to protonate p-nitroaniline (pKa(p-nitroanilinum) ¼ 6.2 in MeCN), indicating that the dOH sites on SZO are rather weak acids. Uniform Brønsted acid strength of dOH sites is fairly common in crystalline solid oxides,113 and this study showed that dOH sites on SZO behave similarly. Correlating Brønsted acidity with a pKa value, which would relate Brønsted acidity of dOH sites on oxides to Brønsted acids in solution, is challenging. Adsorption energies of reactions of oxides containing Brønsted acid sites with molecular bases (pyridine, amines, etc.) do not correlate with the pKa or Hammett acidity (H0) of the protonated bases in solution.114 This is not too surprising because solvation energies associated with placing both the base and the protonated base in solution are often comparable to the energy of the protonation reaction.113 In contrast, plots of DHads and the proton affinity (PA) of the base are linear. A thermodynamic cycle that is consistent with this behavior for the protonation of phosphines on SZO is shown in Fig. 9B, and is closely related to the thermodynamic cycle derived for reactions of zeolites with bases.113,114 The three energies that contribute to the equilibrium adsorption of PR3 to SZO to form [R3PH][SZO] are the DPESZO, defined as the gas phase deprotonate energy of an dOH group on SZO, PAPR3, the proton affinity of the phosphine, and DHion-pair, the energy of ion-paring on the solid oxide. DFT studies are needed to obtain DPESZO. SZO was modeled as periodic (101) or (001) ZrO2 surfaces containing H2SO4 using DFT methods.115 On both surfaces the sulfate engages in tridentate coordination to the ZrO2 surface; the protons are dissociated from the sulfate and lie on a nearby ZrdOdZr bridges (Fig. 10). These results are not entirely surprising because ZrO2 is a Brønsted base. A bidentate binding mode for sulfate is also possible on ZrO2, but molecular dynamics simulations show that bidentate sulfates convert to tridentate sulfates at 800 K, typical crystallization conditions in the synthesis of SZO. The calculated deprotonation energies of these Brønsted acid sites range from 320 to 370 kcal mol−1, significantly higher than the gas phase acidity of H2SO4 (302.3 kcal mol−1),116 but more acidic than simple small molecule silanols (359.3 kcal mol−1).117 Sulfated aluminum oxide (SAO) has similar surface chemistry as SZO. DFT calculations of SAO, using the (110) surface of g-Al2O3, show that H2SO4 adsorbs by reactions with dOH groups on the g-Al2O3 surface to give 3.0 H2SO4 per nm2 and a hydroxyl coverage of 4.4 OH per nm2.118 Similar to SZO, H2SO4 dissociates on the g-Al2O3 surface by protonation of Brønsted basic AldOdAl bridges. On the SAO surface both tridentate sulfate (SA) and bidentate sulfate (SB) sites are stable. Similar models probably apply to other sulfated oxides (e.g. SO4/SnO2, SO4/Fe2O3, SO4/WO3, etc.) and oxides containing polyoxometalate anions.119,120

590

Organometallic Chemistry on Oxide Surfaces

Fig. 10 Equilibrium structure of H2SO4 adsorbed to the (101) surface of ZrO2. Reproduced with permission from reference Haase, F.; Sauer, J. J. Am. Chem. Soc. 1998, 120, 13503–13512.

1.19.3

Reactions of organometallics with oxides by protonolysis

All of the oxides described above contain dOH groups and react with MdR by protonolysis to release RH. There are significant differences in the surface species generated in this reaction that depend on the nature of the support.

1.19.3.1

Reactions of organometallics with partially dehydroxylated silica

Silica adsorbs small molecules functionalized with H-bond acceptors. Recent examples are phosphine functionalized calixarene Ir-clusters that adsorb on silica, and show structural similarities to the clusters in solution.121–124 Detailed solid-state NMR studies of silica contacted with Cp2Fe, which is unreactive toward silanols, or phosphine oxides, which form strong hydrogen bonds with silanols, show that both substrates are dynamic on the NMR timescale,125 suggesting that reorientation or tumbling of adsorbates on silica surfaces occurs more rapidly than the NMR timescale. Similar dynamics across the silica surface are typically not observed in supported organometallics because they are anchored to the surface through ^SidOdM bonds. These reactions are thought to occur through proton transfer or four-center transition states (Fig. 11A). Organometallics containing M]CR2 groups undergo a more complex grafting processes. Ta(]CHtBu)(CHt2Bu)3 reacts with deuterated SiO2-500, which contains 90% ^SidOD, to form ^SiOdTa(]CHtBu)(CHt2Bu)2, neopentane-d1 and neopentane-d0.126 This result indicates that some isolated silanols add to the tantalum alkylidene to form the ^SiOdTa (CHt2Bu)4 intermediates that undergo a-elimination to form ^SiOdTa(]CHtBu)(CHt2Bu)2 (Fig. 11B). This result is in contrast to reactions of tBu3SiOD with Ta(]CHtBu)(CHt2Bu)3 that generate exclusively neopentane-d1,127 showing that simple soluble molecular mimics to isolated silanols do not always behave similarly to silica.128 The addition of silanols to M]CR2 groups is common, and in some cases the product of silanol addition to the silanol is an observable product in the grafting reaction.129,130 Complementary spectroscopic methods are usually required to determine the structure of surface species. In this regard, structural studies of ^SiOdTa(]CHtBu)(CHt2Bu)2 supported on SiO2-700 are emblematic of the rigor achievable using combinations of solid-state NMR spectroscopy and XAS.131 The tantalum LIII edge extended X-ray fine structure (EXAFS) experimental and simulated data are shown in Fig. 12A, and extracted distances are given in Fig. 12B. One Ta]CHtBu scatterer at 1.898(8) A˚ is present, which is close to typical M]CHR bond distances. In addition, the TadO distance is 1.898 A˚ and two TadC distances fit to 2.150(4) A˚ . Finally, this data requires an additional scatter that is the result of the Ta interacting with a nearby SidOdSi bridge. Solid-state NMR spectroscopy is indispensable in studying supported organometallics.43,52 The 13C cross polarization (CP) MAS spectrum of 13C enriched ^SiOdTa(]13CHtBu)(13CHt2Bu)2 contains signals at 247 ppm for the Ta]13CHtBu, 95 ppm for the Tad13CHt2Bu groups, 47 ppm for the TadCH2CMe3/Ta]CHCMe3, and 31 ppm for the TadCH3CMe3/Ta]CHCMe3. The 13C CPMAS NMR assignments are supported by 2D J-resolved experiments that measure 1JC–H couplings, showing that 1JC–H for Ta]13CHtBu is 80 Hz. This value is far below expected values for sp2 hybridized carbons, and is consistent with trends in 1JC–H for supported alkylidenes containing a-agostic structures.14,132 Heteronuclear 2D solid-state NMR experiments often use 13C enriched organometallics, which is obviously cost and labor intensive. DNP NMR spectroscopy gives significant signal enhancements for surface species, offering an alternative to isotopic enrichment,44,54,133 and has been applied directly to organometallics supported on silica.134–136 The presence, or absence, of a M⋯ Od(SiOX) interaction can also be determined using 45Sc NMR spectroscopy. Scandium is quadrupolar and has one high sensitivity NMR active nucleus (45Sc, I ¼ 7/2, g ¼ 6.5081  107 rad T−1 s−1, 100% abundant). In the solid-state, 45Sc NMR spectra are characterized by broad powder patterns that are a result of the interaction of the quadrupole moment of 45Sc with the electric field gradient (EFG) tensor. The quadrupolar coupling constant (CQ) can be extracted from NMR

Organometallic Chemistry on Oxide Surfaces

591

(A)

(B)

Fig. 11 Reaction of a generic organometallic with isolated silanols on SiO2 to by s-bond metathesis (A) or by addition to the alkylidene in Ta(]CHtBu)(CHt2Bu)3 to form ^SiOdTa(]CHtBu)(CHt2Bu)2 after a-elimination (B).

Fig. 12 Characterization data for ^SiOdTa(]CHtBu)(CHt2Bu)2; EXAFS (A) with values for fits to the data (B) and J-resolved solid-state NMR spectroscopy (C). EXAFS and NMR data reproduced with permission from reference Le Roux, E.; Chabanas, M.; Baudouin, A.; de Mallmann, A.; Copéret, C.; Quadrelli, E. A.; Thivolle-Cazat, J.; Basset, J. M.; Lukens, W.; Lesage, A.; Emsley, L.; Sunley, G. J. J. Am. Chem. Soc. 2004, 126, 13391–13399.

spectra of quadrupolar nuclei, and trends in CQ relate to structure at the quadrupolar nucleus.137 For example, the 45Sc{1H} NMR spectrum of the C3 symmetric Sc[N(SiMe3)]3 amide shown in Fig. 13A contains a broad powder pattern that give a CQ of 66.2 MHz.138 Reactions of Sc[N(SiMe3)]3 result in the formation of ^SiOdSc[N(SiMe3)2]2, and 45Sc{1H} NMR studies of this material show that the powder pattern is quantitatively narrower than Sc[N(SiMe3)]3, suggesting that nearby siloxane bridges coordinate to scandium (Fig. 13B). However, the featureless signal precludes accurate simulation of the 45Sc NMR parameters in ^SiOdSc[N(SiMe3)2]2.

592

Organometallic Chemistry on Oxide Surfaces

(A)

(B)

(D)

(C)

Cp* 2Sc–Ph 38.6 MHz Cp* 2Sc–Me 36.2 MHz Cp* 2Sc–Cl 30.0 MHz

Cp* 2Sc–I 27.4 MHz

Cp* 2ScCl(THF) 7.4 MHz

Cp* 2Sc–F 34.2 MHz Cp* 2Sc 30.1 MHz H Cp* 2Sc–Br 29.2 MHz Cp* 2ScF(THF) 26.0 MHz

Cp* 2ScBr(THF) 7.5 MHz Cp* 2ScI(THF) 6.9 MHz

Fig. 13 Static 45Sc{1H} NMR spectrum of Sc[N(SiMe3)3]3 (A) and ^SiOdSc[N(SiMe3)2]2 (B). CQ for Cp 2ScdX and Cp 2ScX(THF) (C) and static 45Sc{1H} NMR spectrum of Cp 2ScdOSi^ (large CQ) and Cp 2Sc(OSi^)(O(SiOX)2) (small CQ) (D). Reproduced with permission from (A and B) reference Vancompernolle, T.; Trivelli, X.; Delevoye, L.; Pourpoint, F.; Gauvin, R. M. Dalton Trans. 2017, 46, 13176–13179 and (C) reference Culver, D. B.; Huynh, W.; Tafazolian, H.; Conley, M. P. Organometallics 2020, 39, 1112–1122.

14-electron organoscandium complexes of the type Cp 2ScdX (Cp ¼ pentamethylcyclopentadienyl, X ¼ halide, alkyl, aryl) were studied in detail to understand s-bond metathesis and olefin insertion reactions.139 The bulky Cp ligands restrict dimerization and provide solubility in hydrocarbon solvents. The CQ of Cp 2ScdR,140 Cp 2ScdX,141 and Cp 2ScX(THF)141 extracted from 45 Sc{1H} NMR spectra are summarized in Fig. 13C. Base-free Cp 2ScdX (X ¼ F, Cl, Br, I) and Cp 2ScdR (R ¼ Me, Et, Ph) have relatively large CQ values >27 MHz that are larger than high-symmetry scandium complexes142 or crystalline porous materials,143 but smaller than Sc[N(SiMe3)]3. In the alkyl series, Cp 2ScdEt has a smaller CQ than Cp 2ScdMe because the former contains a b-agostic structure.140 In the halide series, base-free Cp 2ScdX have significantly larger CQ than Cp 2ScX(THF) (X ¼ F, Cl, Br, I) because the THF adducts have higher symmetry at Sc compared to the base-free halides, which results in some degree of p-bonding from the halide to scandium.141 The reaction of Cp 2ScdMe with SiO2-700 forms Cp [email protected] The 45Sc{1H} NMR spectrum of Cp 2Sc@SiO2 contains a broad feature that simulates to two different scandium environments giving CQ of 35.4 and 21.9 MHz, respectively (Fig. 13D). This result indicates that both Cp 2ScdOSi^ (large CQ) and Cp 2Sc(OSi^)(O(SiOX)2) (small CQ) are present in Cp 2Sc@SiO2. This interpretation was supported with studies of molecular Cp 2ScdOR (R ¼ CMe2CF3, CMe(CF3)2, C(CF3)3, SiPh3) that are characterized by large CQ values (29.2–35.1 MHz), and DFT modeling showing that Cp 2Sc(k1-OSi(OMe)3) has a much larger CQ than Cp 2Sc(k2-OSi(OMe)3).

Organometallic Chemistry on Oxide Surfaces

593

(A)

(B)

(C)

Fig. 14 Homoleptic organolanthanides supported on SiO2 containing secondary interactions before and after grafting.

The low-coordinate organolanthanides shown in Fig. 14 are interesting probes of how the silica surface affects coordination environment because lanthanides typically have high coordination numbers (>6), and each of these complexes contain additional secondary interactions between the organic fragment and the lanthanide. Ln[CH(SiMe3)2]3 (Ln ¼ Y,144 La,145 Ce,144 Sm,145 Lu130) contain secondary 3c-2e m-Me interactions between proximal SidMe groups and the lanthanide.130 The reaction of Lu[CH (SiMe3)2]3 with SiO2-700 forms ^SiOdLu[CH(SiMe3)2]2, which coordinates to a siloxane bridge and contains one Lu ⋯ MedSi interaction at 2.80(2) A˚ , assigned to the 3c-2e m-Me, from EXAFS data (Fig. 14A).130 The Lu ⋯ MedSi distance in ^SiOdLu [CH(SiMe3)2]2 is significantly shorter than the three Lu ⋯MedSi distances in Lu[CH(SiMe3)2]2 (2.973(3) A˚ ), but longer than the Lu ⋯ MedSi distances in Lu[CH(SiMe3)2]2[O-2,6-tBu2-C6H3] (2.69  0.02 A˚ ). Related observations were made when grafting organolanthanides containing Ln⋯ HdSi secondary interactions. Y[N(tBu) (SiHMe2)]3 contains three Y ⋯ HdSi secondary interactions that have characteristic 1JSiH values. Reactions of Y[N(tBu) (SiHMe2)]3 and partially dehydroxylated silica forms ^SiOdY[N(tBu)(SiHMe2)]2 (Fig. 14B).146 Solid-state NMR studies show that 1JSiH values in ^SiOdY[N(tBu)(SiHMe2)]2 are significantly higher values than the Y[N(tBu)(SiHMe2)]3, suggesting that the Y ⋯HdSi interaction is weaker in the supported yttrium species, likely due to coordination of a siloxane bridge to yttrium. Reactions of La[CH(SiHMe)3]3,147 which contains six La ⋯ HdSi interactions (1JSiH ¼ 137 Hz), with partially dehydroxylated silica forms ^SiOdLa[CH(SiHMe)3]2 (Fig. 14C). J-resolved 1Hd29Si NMR spectroscopy of ^SiOdLa[CH(SiHMe)3]2 gives 1JSiH of 167 Hz, consistent with two La ⋯ HdSi interactions, again suggesting that siloxane bridges displace La ⋯ HdSi secondary interactions.148 These examples show the impact of silica as a ligand, which is generally but not exclusively consistent with the silica surface acting as an LX type ligand (X ¼ ^SiOd; L ¼ SidOdSi) in the Green formalism.149 Viewed through the lens of the secondary interactions shown for the low-coordinate lanthanides in Fig. 14, the L-type coordination of silica to the supported lanthanides appears to weaken Y⋯ HdSi and La⋯ HdSi interactions, which is probably due to a combination of coordination of siloxane bridges to the lanthanide and/or dynamics on the NMR timescale that exchange Ln ⋯ HdSi with uncoordinated HdSi present in the ligand structure. However, the siloxy X-type ligand increases the Lewis acidity of the lanthanide. This is most evident in ^SiOdLu[CH(SiMe3)2]2, which has a much shorter Lu⋯ MedSi distance than Lu[CH(SiMe3)2]3. The silica surface is amorphous, and the presence of L-type interactions is not necessarily uniform across all species supported on silica. This was evident in the 45Sc NMR studies of Cp 2Scdsites supported on silica showing that both Cp 2Sc(k1-OSiOX) and Cp 2Sc(k2-OSiOX) are products in grafting reactions of Cp 2ScdMe and SiO2-700.77 Similar heterogeneities in surface speciation were also encountered in 27Al NMR studies of Al[N(SiMe3)3]2Cl(THF) supported on silica.150 This implies that not all silanols on the surface of partially dehydroxylated silica are in identical chemical/coordination environments.151

594

Organometallic Chemistry on Oxide Surfaces

Fig. 15 Heterogeneity of silanols on silica from reactivity of Os(]CHtBu)2(CHt2Bu)2 or [AuN(SiMe3)2]4.

Indeed, there are cases where the heterogeneity of the silica surface is much more pronounced. The reaction of Os(] CHtBu)2(CHt2Bu)2152 with SiO2-700 results in only 30% conversion of silanols, independent of reaction time or number of equivalents of Os added to the reaction mixture, to form the see-saw ^SiOdOs(^CtBu)(CHt2Bu)2 (Fig. 15), and ^SiOdOs(] CHtBu)2(CHt2Bu) expected from protonolysis of one OsdCHt2Bu does not form in this reaction.153 The grafting reaction was studied using DFT methods, which showed that the isomerization of the tetrahedral osmium alkylidene to the see-saw osmium alkylidyne is favored by the weak s-donor siloxide. The silanols that do not react with Os(]CHtBu)2(CHt2Bu)2 probably cannot accommodate the transition state necessary for isomerization to the alkylidyne product. Another example is the reaction of SiO2-700 and [AuN(SiMe3)3]4 followed by contacting with excess PMe3 forms ^SiOdAuPMe3, the expected product, and small amount of [(Me3P)2Au][OSi^] ion pairs (Fig. 15). This reaction is quite unusual because the silanols on silica normally form ^SidOdM species rather than forming ion-pairs. The origin of this behavior is not clear, but is may be related to the amorphous silica surface that promote some ^SidOdM sites to ionize. Silicas dehydroxylated at lower temperatures are expected to result in more complicated surface speciation of supported organometallics because a complex network of isolated an hydrogen-bonded silanols are present on these materials. The reaction of W(]NAr)(]CHtBu)(CHt2Bu)2 (Ar ¼ 2,6-diisopropylphenyl) with SiO2-200 generates neopentane and the three different supported W-species shown in Fig. 16, and exemplifies the complexity that can be encountered with silicas dehydroxylated at temperatures lower than 700  C.154 Notably absent from this mixture is the tris-grafted W species. However, this complexity is dependent on the organometallic used in grafting reactions. Zr(CHt2Bu)4, Ta(]CHtBu)(CHt2Bu)3, M(]O)(CH2EMe3) (E ¼ C, Si), and [Fe(Mes)2]2 react with SiO2-200 to form well-defined siloxy surface species.70,155–160 Supporting an organometallic complex on silica can have favorable effects on catalytic reactions. This is a common feature in SOMC, and has been thoroughly reviewed.36–42 The impact of a siloxy ligand on catalytic behavior of is most thoroughly studied for supported Group VI alkylidenes in olefin metathesis, and will be briefly covered here to show these favorable effects.161,162 Olefin metathesis with supported catalysts is mechanistically similar to homogeneous olefin metathesis reactions, shown for a generic metal-imido-alkylidene in Fig. 17.72 Coordination of an olefin to a 14-electron alkylidene, followed by 2+2 cycloaddition forms the trigonal bipyramidal (TBP) metallacyclobutane. The TBP isomer is on the metathesis catalytic cycle, but can isomerize to the square pyramidal (SP) isomerwhich is not on the metathesis catalytic cycle. These isomers have very different NMR chemical shift patterns, which relates to metathesis activity.163 Supported metathesis catalysts generally have higher activities than related homogeneous catalysts, and quantitative models that predict metathesis activity as a function of ligand environment for supported imido-alkylidene catalysts are available.164,165 There are two key reasons for the increased activity of supported metathesis catalysts. First, the dissymmetric ligand environment at the metal center reduces the barrier to olefin coordination and 2 +2 cycloaddition,72 a key design strategy in state-of-the-art homogeneous catalysts for Z-selective olefin metathesis.166–168 Second, the siloxy ligand is small, providing access to unique steric environments that are inaccessible in solution because small alkylidenes decompose by bimolecular decomposition.169,170 For example, (AdPO)2W(]CHtBu)(]O) contains very bulky 2,6-diadamantylphenoxide ligands, and only reacts with olefins in the presence of B(C6F5)3 to form W(IV)-olefin complexes.171 Supporting (AdPO)2W(]CHtBu)(]O) on SiO2-700 results in formation of AdPOH and ^SiOdW(]CHtBu)(]O)(OPdA), a very active supported metathesis catalyst that mediates >75,000 turnovers in 1-hexene metathesis (Fig. 18A).18

Organometallic Chemistry on Oxide Surfaces

595

Fig. 16 Reactivity of SiO2-200 in a representative grafting reaction.

Fig. 17 Mechanism of olefin metathesis.

Treating supported organometallics with H2 is a well-known strategy to generate low-coordinate metal hydride species not typically accessible in solution (vide infra). This is a result of the beneficial site isolation provided by the siloxy ligand that prevents aggregation, similar to deactivation pathways for alkylidenes mentioned above. For example, the reaction of (tBuCH2)3Ta] Ir(H)2Cp with silica by loss of neopentane forms well-defined species that react with H2 to form ^SiOdTa(CHt2Bu)(H)(] Ir(H)2Cp ).172 This is in contrast to reactions of (tBuCH2)3Ta]Ir(H)2Cp with H2 in solution that forms tetranuclear clusters (Fig. 18B).

1.19.3.2

Reactions of organometallics with dOH groups present on sulfated oxide surfaces

SMOs also react with organometallics by protonolysis, and are relatively recent supports for SOMC. The dOH sites on SMOs are more acidic than silanols on silica, which should affect the metal⋯ support interaction. The reaction of Cp 2ZrMe2 oxides is a representative example showing this effect (Fig. 19).173 SiO2 reacts with Cp 2ZrMe2 to form Cp 2Zr(Me)(OSi^) and methane. The 13 C CPMAS spectrum of Cp 2Zr(Me)(OSi^) contains a signal at 31.5 ppm assigned to the ZrdMe, close to values obtained for Cp 2ZrMe(OR).174 Consistent with this assignment, Cp 2Zr(Me)(OSi^) does not catalyze olefin polymerization reactions because the siloxy ligand does not behave as a weakly coordinating anion. However, SAO reacts with Cp 2ZrMe2 to form methane and [Cp 2ZrMe][SAO] ion-pairs. This assignment is consistent with the 13C CPMAS NMR signal for the ZrdMe+ (d ¼ 46 ppm), which is close to the 13C NMR chemical shift for [Cp 2ZrMe][MeB(C6F5)3] (d ¼ 50.4 ppm).175 Unlike Cp 2Zr(Me)(OSi^), [Cp 2ZrMe] [SAO] is very active in ethylene polymerization reactions, establishing the weakly coordinating behavior of sulfates on SAO. A fundamental question in reactions of SMOs with organometallics is if the metal coordinates to a MdOdM bridge, the original location of the acidic proton, or to sulfate on the oxide surface. DFT studies of [Cp2ZrMe]+ interacting with deprotonated SAO show that only SA or SB interact with the electrophilic organozirconium fragment (Fig. 20). AldOdAl bridges do not interact with [Cp2ZrMe]+ because of unfavorable steric interactions between the sulfated alumina surface and the organozirconium fragment. Importantly, the ZrdO distances modeled from DFT (ZrdOSA ¼ 2.24 A; ZrdOSB ¼ 2.42 A) are close to those obtained experimentally from EXAFS measurements (ZrdO ¼ 2.37 A).118

596

Organometallic Chemistry on Oxide Surfaces

(A)

(B)

Fig. 18 Effects of a siloxy ligand from silica on the catalytic properties of a W-oxo alkylidene (A) and the formation of reactive Ta-H species (B).

Fig. 19 Effect of the oxide support on surface speciation.

Organometallic Chemistry on Oxide Surfaces

597

Fig. 20 DFT optimized structures of Cp2ZrMe+ interacting with SA (left) or SB (right) on sulfated aluminum oxide. Reproduced with permission from reference Williams, L. A.; Guo, N.; Motta, A.; Delferro, M.; Fragala, I. L.; Miller, J. T.; Marks, T. J. Proc. Natl. Acad. Sci. U.S.A. 2013, 110, 413–418.

Examples of supported [M][SMO] ion pairs are shown in Fig. 21. The reader may recognize the emphasis on supported species that contain similarities with well-known Group 426,57 or Group 10176 organometallics for olefin polymerization reactions. Indeed, organozirconium and organohafnium complexes supported on SMOs behave similarly to electrophilic ion-pairs containing weakly-coordinating anions (e.g. MeB(C6F5)3, B(C5F5)4),173,177–184 a property that also applies to Ni and Pd complexes containing bulky a-diimine ligands.185,186 The applications of SMOs as supports for well-defined organometallics is not limited to olefin polymerization reactions. Cp IrMe2(PMe)3 reacts with SMOs to form [Cp IrMe(PMe3)][SMO] that are active in H/D exchange reactions,187 and (dmPhebox)Ir(OAc)2 reacts with SZO to form electrophilic Ir sites for CdH bond activation an olefin hydrogenation reactions.188 SZO also reacts with allyltriisopropylsilane to formation of [iPr3Si][SZO] ion pairs, a clear demonstration of the weakly coordinating behavior of the sulfates on SZO.189 R3Si+ ions were the subjects of intense academic debate in the mid1990s190–193 that was not settled until sterically bulky Mes3Si+ was isolated and crystallographically characterized as the [CHB11Me5Br6] salt.194 R3Si+ ions are very sensitive to the nature of the weakly coordinating anion,195,196 show characteristic deshielded 29Si NMR chemical shifts,197 and are known to activate CdF bonds.198 The solid-state 29Si NMR spectrum of [iPr3Si] [SZO] contains a major signal at 53 ppm,  39 ppm downfield from ^SidOdSiR3,199–201 consistent with a Lewis acidic iPr3Si+ fragment in this material. [iPr3Si][SZO] activates sp3 CdF bonds in trifluorotoluene, perfluorotoluene, 1-adamantyl fluoride, and 1H,1H,2H-perfluorohexene in the presence of excess HSiEt3 to give hydrocarbon products. [iPr3Si][SZO] is more reactive in hydrodefluorination reactions than high surface area AlF3,202 the only other heterogeneous catalyst for this reaction, but is significantly less reactive that R3Si+ ions containing weakly coordinating borate or carborane anions.203–205 The most common reaction pathway of a SMO with an organometallic is protonolysis of the organometallic ligand, though at high dehydroxylation temperatures side reactions can be observed due to formation of pyrosulfate sites on the SMO surface. This was encountered in reactions of Cp IrMe2(PMe3) with SAO dehydroxylated at 500  C (Fig. 22).206 Performing this reaction in C6D6 resulted in deuterium incorporation into the methyl groups of the Cp ligand, which was not expected for a simple protonolysis reaction. This behavior is a result of a formal H-atom transfer grafting process, shown in Fig. 16, that forms the Ir-tetramethylfulvene as an intermediate. Consistent with this model, DFT studies showed that electron transfer to a pyrosulfate on SAO is favorable by −21.5 kcal mol−1, and that similar electron transfer to a monosulfate on SAO is unfavorable by +17.6 kcal mol−1. Pyrosulfates can be avoided by lowering sulfate loading and treating SAO at lower dehydroxylation temperatures.206

1.19.3.3

Factors affecting formation of ^SidOdM or [M][oxide]

Partially dehydroxylated SiO2 and SMOs are representative oxides that react with organometallics, by protonolysis of MdR groups by dOH groups on the oxide surface. However, these two oxides give very different reaction products: SiO2 forms ^SidOdM and SMOs form [M][SMO]. Understanding how to control this speciation has obvious implications for catalysis. For example, reactions of Cp IrMe2(PMe3) with dOH sites form well-defined species for H/D exchange reactions (Fig. 23).187 Silica supported Ir species are unreactive in H/D exchange, but SMOs are active in this reaction. Ir species supported on ZrO2 treated with boric acid fall somewhere between these two extremes. This trend is due to the ability of the Ir fragment to dissociate from the anion-doped supports to form transient [Cp IrMe(PMe3)]+ that contain an open coordination site at iridium to activate substrate in the H/D exchange reaction. Acidity trends are useful to consider when evaluating such reactivity trends because anions (X−) form weaker ion pairs as the conjugate acid (HX) becomes stronger. Boric acid is a much weaker acid than sulfuric acid. Therefore, oxides doped with boric acid should form stronger ion-pairs than oxides doped with sulfate anions. Fig. 24 shows gas phase acidity (HX ! H+ + X−) for H3SidOH, triflic acid, and H[CH6B11Cl6]. As the anion becomes more weakly coordinating the energy of the gas phase acidity reaction decreases, as expected. This property also relates to ion-pairing207 and to the 29Si NMR chemical shift for R3SidX197; selected 29Si NMR chemical shifts for selected R3SidX and [R3Si][X] are also shown in Fig. 24. This suggests that a combined understanding of gas phase acidity properties of oxides and measurement of 29Si NMR chemical shifts for alkylsilane functionalized oxides will provide information about formation of ^SidOdSiR3 or [R3Si][Oxide], as was seen in triakylsilyl-functionalized silica and [iPr3Si][SZO].

598

Organometallic Chemistry on Oxide Surfaces

Fig. 21 Well-defined species supported on SMOs.

(A)

(B)

Fig. 22 Reactivity of pyrosulfates with Cp IrMe2(PMe3).

Organometallic Chemistry on Oxide Surfaces

599

Fig. 23 H/D exchange of arenes and alkanes catalyzed by supported Cp IrMe(PMe3) on oxides doped with different anions. Reproduced with permission from reference Witzke, R. J.; Chapovetsky, A.; Conley, M. P.; Kaphan, D. M.; Delferro, M. ACS Catal. 2020, 11822–11840, https://doi.org/10.1021/acscatal.0c03350.

Fig. 24 Relationship of gas phase acidity for HX and 29Si NMR chemical shift of R3SidX.

Generation of a surface more weakly coordinating than SZO would require more acidic dOH sites and a more deshielded 29Si NMR chemical shifts in a trialkylsilyl-functionalized material. Lewis acids react with a Bronsted acids to form very strong Bronsted acids in solution.208 Similarly, Al(OC(CF3)3)3(PhF), a very strong Lewis acid,209 reacts with the weak Bronsted acid silanols on SiO2F F 76 The bridging silanols have a calculated gas 700 to form well-defined bridging silanols (^SidOHdAl(OR )3, R ¼ C(CF3)3). −1 phase acidity of 267.2 kcal mol , significantly more acidic than triflic acid and dOH sites on SZO. Contacting ^SidOHdAl (ORF)3 with allyltriisopropylsilane results in formation of [iPr3Si][RFO3AldOSi^], which contains a 29Si CPMAS NMR chemical shift at 70 ppm. These values indicate that the anions in [RFO3AldOSi^] are more weakly coordinating than SZO, but more strongly coordinating than carborane anions. Similar strategies may result in more weakly coordinating surfaces, but requires development of neutral Lewis acids stronger than Al(OC(CF3)3)3(PhF).

1.19.4

Reactivity of strained ^SidOdSi^ bridges on dehydroxylated silica surfaces

The electrophilic grafting pathway shown in Fig. 2 is presently limited to silica. The simplified dehydroxylation model describing Y of surface silanols as a function of temperature shown in Fig. 4 implies that as dehydroxylation temperatures increase, the

600

Organometallic Chemistry on Oxide Surfaces

(A)

(B)

Fig. 25 Opening of siloxane bridges with NH3 (A). Restructuring of silica thru reactions of geminal silanols capped with dSnBu3 groups (B).

population of strained siloxane bridges also increases. Indeed, strained ^SidOdSi^ bridges are often suggested to be reactive in reactions of silica partially dehydroxylated at very high temperatures (>700  C). For example, strained siloxane bridges in SiO2-1200 reacts with NH3 to form ^SidNH2 and ^SidOH pairs that are in close proximity (Fig. 25A).210,211 Restructuring of silica surfaces was also encountered in grafting reactions with Bu3Sn(allyl), and has a pronounced dependence on the temperature of dehydroxylation.212 SiO2-700 reacts with Bu3Sn(allyl) to form ^SiOdSnBu3 (major, d117Sn ¼ 122 ppm), (^SiO)2SnBu2 (minor, d 119Sn ¼ −8 ppm), propene, and butane. The amount of butane evolved in this reaction, a product in the formation of (^SiO)2SnBu2, increases as silica dehydroxylation temperature increases. These results were rationalized by reactions of geminal ]Si(OH)2 with Bu3Sn(allyl) to form ]Si(OH)(OSnBu3) that subsequently react with strained siloxane bridges to form (^SiO)2SnBu2 and butane, shown in Fig. 25B. Consistent with this model, SiO2 dehydroxylated between 200 and 500  C reacts with Bu3Sn(allyl) to form only ^SiOdSnBu3 and propene, because significant quantities of strained siloxane bridges are not present on silica treated at these temperatures. More dramatic silica surface reorganization through opening of siloxane bridges is common in grafting reactions with organoaluminum213–220 and organogallium reagents.221–223 For example, AliBu3 reacts with SiO2-500 in pentane to form a mixture of surface products shown in Fig. 26. These products form as a result of protonolysis of AldiBu groups with silanols to form unobserved ^SiOdAliBu2 sites that open siloxane bridges by transferring iBu groups to silicon and forming new AldO bonds (Fig. 26A). This reaction forms ^SidiBu, ]Si(iBu)2 and dSi(iBu)3 from 29Si CPMAS NMR experiments. In addition to the organoaluminum product shown in Fig. 26A, aluminum sites lacking AldiBu groups are also present from 27Al NMR studies. In contrast, AliBu3 reacts with SiO2 in diethyl ether to form the monomeric (^SiO)2AliBu(OEt2).218,219 This product also forms from a combination of protonolysis of AldiBu and alkyl transfer to form ^SidiBu, but does not form alkylsilane surface sites, suggesting that Et2O inhibits further reaction of organoaluminum sites with siloxane bridges. The divergent reactivity of AliBu3 with silica as a function of solvent shows how subtle changes in reaction conditions can result in very different surface species. High oxidation state transition metals also open siloxane bridges during low temperature grafting reactions. CrO2Cl2 reacts with SiO2-800 to form ^SiOdCrO2Cl and ^SiCl sites,224,225 and Re2O7 reacts with SiO2-1000 to form ^SiOdReO3.226 TiCl4 reacts with SiO2 to form ^SiOTiCl3 and HCl, and thermal annealing results in surface restructuring to form (^SiO)2TiCl2 with release of TiCl4.227 This reactivity is rare for organometallics because protonolysis of MdR groups by silanols is the most common reaction pathway, as discussed above. However, highly reactive La(CH2Ph)23THF grafts onto SiO2 to form ^SiOdLa(CH2Ph)2(THF)x, (^SiO)2La(CH2Ph)(THF)x, and ^SidCH2Ph.228 Siloxane bridge opening is more common in post-synthetic reaction chemistry of a supported organometallic. Hydrogenolysis of ^SiOdTa(]CHtBu)(CHt2Bu)2 or ^SiOdZr(CHt2Bu)3 at 150  C results in formation of equilibrium mixtures of (^SiO)2TaH3 and (^SiO)2TaH229 or (^SiO)3ZrdH/(^SiO)2ZrH2 (Fig. 27A).230 This reaction involves formation of ^SiOdMHx that react with siloxane bridges by hydride transfer to form ^SiH sites and new ^SiOdM bonds. These very reactive MdH surface sites reduce N2 to form metal-imido-amide products231 and also engage in s-bond metathesis, b-alkyl transfer, and/or a-elimination that can achieve catalytic cycles for alkane hydrogenolysis or alkane metathesis.27 In the absence of H2, thermolysis of ^SiOdTaMe4 results in formation of methane, ^SiOdTa(]CH2)Me2, and (^SiO)2Ta(]CH2)Me; the latter of which forms by successive a-elimination and opening of a nearby siloxane bridge (Fig. 27B).232 Similar thermolytic chemistry was also observed in ^SiOdWMe5 (Fig. 27C).233

Organometallic Chemistry on Oxide Surfaces

601

Fig. 26 Reaction of AliBu3 with SiO2 in pentane or ether gives different surface species.

Thermal decomposition of M[OSi(OtBu)3]n supported on silica also results in reorganization of the silica surface. Originally reported by Tilley and co-workers, calcination of M[OSi(OtBu)3]n results in the release of isobutene and isobutanol and formation of well-defined MOX/SiOx materials.30,234,235 M[OSi(OtBu)3]n supported on SiO2 form high oxidation state (^SiO)n-1MOH when calcined (Fig. 27D), and are active in oxidation reactions.236–241 This chemistry is also directly applicable to the thermal decomposition of M[OSi(OtBu)3]n supported on SiO2 with conservation of the oxidation state of the starting material if thermolysis is conducted under vaccum.51 The thermolytic precursors for this chemistry are readily available through reactions of MdNR2 with HOSi(OtBu)3 or through salt-metathesis, and M[OSi(OtBu)3]n react with silica through protonolysis with dOH sites on oxides. During thermolysis the dOtBu fragments decompose to form isobutene and dOH groups that undergo condensation with nearby siloxane bridges. Though thermolysis usually requires 350  C, the coordination environment of metal sites formed by heat treatment are usually quite similar to the coordination environment of the M[OSi(OtBu)3]n starting material.51 Similar siloxane bridge reorganization is also known for metal amides supported on silica after thermal annealing.242,243

1.19.5

Alkyl abstraction by Lewis sites on Al2O3

Much of the early work in SOMC by Ballard244–246 and Yermakov247,248 was devoted to determining the active sites in Ziegler-Natta type olefin polymerization catalysts, usually by supporting ZrR4 onto silica or alumina (Ballard) followed by H2 treatment (Yermakov). These are pioneering examples of SOMC, but the supported species were far less active than typical Ziegler-Natta catalysts. One important result from these studies was that Zr(CH2Ph)4 supported on alumina was at least 100 times more active than ZrR4 supported on silica.244 Several years later, Burwell and Marks showed that Cp 2ThMe2 and Cp 2UMe2 supported on alumina were also quite active in polymerization of ethylene.249,250 Alumina dehydroxylated at 1000  C (Al2O3-1000) does not contain dOH sites (Fig. 6C), and must contain exposed Lewis sites on the surface of the oxide. The 13C CPMAS spectrum of Al2O3-1000 contacted with Cp 2Th( Me)2 contains a signal at 71 ppm, 2.6 ppm downfield from Cp 2Th( Me)2, and a signal at −20 ppm assigned to  MedAlOX, consistent with the assignment of [Cp 2Th Me][ MedAlOX] ion-pairs as the product of this reaction (Fig. 28).251 The modest downfield shift for the thorium methyl in [Cp 2ThdMe][MedAlOX] is consistent with systematic studies of the reaction of Cp 2ThMe2 with [Et3NH][X] to form [Cp 2ThMe][X] (X ¼ BPh4; d ¼ 71.8, B(C6F5)4; d ¼ 78.1) or [Cp 2ThMe(THF)2][BPh4] (d ¼ 69.3).252 The 13C NMR chemical shift for the ThdMe in Cp 2ThMe(OSiMet2Bu) or Cp 2ThMe(OSi^), 59.2 and 59.0 ppm respectively, is also consistent with the assignment of the ion pair in [Cp 2Th Me][ MedAlOX].24 To the best of our knowledge this is the first characterized example of alkyl abstraction by an oxide to form ion-pairs, and to place this result in context, this study was reported before isolation of Cp2ZrMe(L)+ cations253,254 or related base-free derivatives.255

602

Organometallic Chemistry on Oxide Surfaces

(A)

(B) (C)

(D)

Fig. 27 H2 treatment of organometallics on silica (A). Thermolysis of ^SiOdTaMe4 (B) and ^SiOdWMe5 (C). Thermolytic molecular precursors supported on SiO2 (D). All processes involve electrophilic opening of siloxane bridges.

As discussed above, g-Al2O3-1000 contains phase impurities because fully dehydrated g-Al2O3 is not stable. g-Al2O3-500 contains both dOH sites and Lewis sites. The reaction of Cp Zr(13CH3)3 with Al2O3-500 results in the formation of methane, indicating that dOH sites on alumina react with ZrdMe groups, and the 13C CPMAS spectrum contains signals at 42, 35, and −11 assigned to the structures shown in Fig. 29.256 The mixture of surface sites encountered with partially dehydroxylated g-Al2O3 is common for d0 metals, and depends on the structure of the organometallic complex.257–261 An exception to this trend is the reaction of W(^CtBu)(CH2tBu)3 with Al2O3-500 that [(OS)3AlIVO]W(^CtBu)(CH2tBu)2 as a major species. In contrast to the other d0 organometallics, the alkyl and alkylidyne fragments interacting with nearby AlSOH groups according to IR and NMR spectroscopies, but do not interact to an appreciable degree with Lewis sites.257,262

Organometallic Chemistry on Oxide Surfaces

603

Fig. 28 Alkyl abstraction by Al2O3-1000.

Fig. 29 Surface species formed in reactions of Al2O3-500 with organometallics.

1.19.6

Heterolytic activation of CdH bonds on Al2O3

Heterolytic cleavage of a CdH bond across a ^EdOdE^ bridge is very common for reactions of small molecules with metal oxides.263 This reaction pathway is important to generate key organometallic intermediates with well-defined metal sites supported on oxides for olefin polymerization and alkane dehydrogenation reactions.51 However, this reaction pathway is the least common for organometallics interacting with oxides. In fact, the only well-characterized example is the reaction of MeReO3 with g-Al2O3, which proceeds by heterolytic cleavage to form ^AldCH2ReO3 (10%) and adsorbed MeReO3 (90%, Fig. 30), a mixture that is active in olefin metathesis at room temperature.264 The 13C CPMAS NMR spectrum of this mixture contains signals at 30 ppm, assigned to adsorbed MeReO3, and at 66 ppm assigned to ^AldCH2ReO3. Isotopic exchange with 13C-labelled ethylene shows that only ^AldCH2ReO3 is active in olefin metathesis reactions and shows that adsorbed MeReO3 is not active in olefin metathesis. DFT analysis of this reaction shows that Re]O sites coordinate to Lewis sites on Al2O3, and that when the m-methylene groups are bound to AlIII sites the barrier to form active Re]CH2 species is sufficiently low to account for the room temperature activity of MeReO3/Al2O3 in olefin metathesis reactions.265 This model also relates to the Re2O7/Me4Sn/Al2O3 catalyst that is the only industrially relevant olefin metathesis catalyst that is active at room temperature.266 Though the alkylidene in the MeReO3/ Al2O3 catalyst has thus far eluded detection, key metallacycle intermediate on the metathesis catalytic cycle (Fig. 17) were characterized by solid-state NMR spectroscopy.265 MeReO3 also reacts with amorphous a-Al2O3 and g-Al2O3 contacted with CCl4 to form sites that are more reactive and stable than MeReO3/Al2O3, though the grafting pathway appears to favor adsorption of MeReO3 onto Lewis sites instead of heterolytic cleavage of the CdH bond.267,268

604

Organometallic Chemistry on Oxide Surfaces

Fig. 30 Reaction of MeReO3 with Al2O3 to form adsorbed and m-methylene species, the latter of which forms unobserved alkylidenes in the metathesis of olefins. Reproduced from reference Valla, M.; Wischert, R.; Comas-Vives, A.; Conley, M. P.; Verel, R.; Copéret, C.; Sautet, P. J. Am. Chem. Soc. 2016, 138, 6774–6785.

1.19.7

Conclusion

In our opinion, a key issue in designing catalytically active sites on surfaces is controlling the how the active site interacts with the support. This article described three common oxides and their behaviors toward organometallics to form well-defined active sites that are tethered to the surface through covalent bonds, as in ^SidOdM supported on silica, or electrostatic ion-pairs, as in [M] [SMO] from protonolysis reactions or [M][MedAlOX] from alkyl abstraction reactions. Applications of basic principles familiar to chemists can be used to predict the products of organometallic grafting reactions. The key factors that affect the products of grafting reactions are acidity of dOH groups and the presence or absence of Lewis sites. dOH groups on SMOs are more acidic than silanols on silica, which explains the difference in the products of organometallic grafting reactions using these two supports. Similarly, Al2O3 contains Lewis sites that can abstract alkyl groups from organometallics, and is similar to reactions of common strong borane Lewis acids with organometallics in solution. Reactions by opening of strained ^EdOdE^ bridges, which is predominantly limited to grafting on SiO2, are less predictable, but generally apply when very electrophilic surface species are formed as reaction products. This reaction transfers alkyl groups from ^EOdMR to the surface to form (^EO)2M sites and RdE^. The least common reaction pathway is grafting by heterolytic cleavage of a CdH bond across ^EdOdE^ bridges, but is mechanistically related to reactions of highly dehydroxylated oxides with alkanes or water. Only three oxides are discussed in this article, because there are few other oxides with a sufficient number of thoroughly characterized supported organometallics.50 As the field continues to grow we expect that other oxides will become important “ligands,” either as spectators or as participants in catalytic reactions, for well-defined organometallics in SOMC.

Acknowledgment The authors thank the National Science Foundation for support (CHE-1800561).

References 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20.

Koval, C. A.; Lercher, J.; Scott, S. L. Basic Research Needs for Catalysis Science to Transform Energy Technolgies; Gaithersburg: Maryland, 2017. Handbook of Heterogeneous Catalysis, 1st Ed.; Wiley-VCH, 1997. Munnik, P.; de Jongh, P. E.; de Jong, K. P. Chem. Rev. 2015, 115, 6687–6718. Schwarz, J. A.; Contescu, C.; Contescu, A. Chem. Rev. 1995, 95, 477–510. Eskandari, S.; Tate, G.; Leaphart, N. R.; Regalbuto, J. R. ACS Catal. 2018, 8, 10383–10391. Serp, P.; Kalck, P.; Feurer, R. Chem. Rev. 2002, 102, 3085–3128. Schaidle, J. A.; Habas, S. E.; Baddour, F. G.; Farberow, C. A.; Ruddy, D. A.; Hensley, J. E.; Brutchey, R. L.; Malmstadt, N.; Robota, H. Catalysis; The Royal Society of Chemistry, 2017; vol. 29; pp 213–281. Banks, R. L.; Bailey, G. C. Ind. Eng. Chem. Prod. Res. Dev. 1964, 3, 170–173. Ivin, K. J.; Mol, J. C. Olefin Metathesis and Metathesis Polymerization; Elsevier Science, 1997. Lwin, S.; Wachs, I. E. ACS Catal. 2014, 4, 2505–2520. Schrock, R. R. Angew. Chem. Int. Ed. 2006, 45, 3748–3759. Grubbs, R. H. Angew. Chem. Int. Ed. 2006, 45, 3760–3765. Peryshkov, D. V.; Schrock, R. R.; Takase, M. K.; Müller, P.; Hoveyda, A. H. J. Am. Chem. Soc. 2011, 133, 20754–20757. Blanc, F.; Basset, J. M.; Copéret, C.; Sinha, A.; Tonzetich, Z. J.; Schrock, R. R.; Solans-Monfort, X.; Clot, E.; Eisenstein, O.; Lesage, A.; Emsley, L. J. Am. Chem. Soc. 2008, 130, 5886–5900. Petroff Saint-Arroman, R.; Chabanas, M.; Baudouin, A.; Copéret, C.; Basset, J. H.; Lesage, A.; Emsley, L. J. Am. Chem. Soc. 2001, 123, 3820–3821. Le Roux, E.; Taoufik, M.; Chabanas, M.; Alcor, D.; Baudouin, A.; Copéret, C.; Thivolle-Cazat, J.; Basset, J. M.; Lesage, A.; Hediger, S.; Emsley, L. Organometallics 2005, 24, 4274–4279. Rataboul, F.; Chabanas, M.; De Mallmann, A.; Copéret, C.; Thivolle-Cazat, J.; Basset, J.-M. Chem. A Eur. J. 2003, 9, 1426. Conley, M. P.; Forrest, W. P.; Mougel, V.; Coperet, C.; Schrock, R. R. Angew. Chem. Int. Ed. 2014, 53, 14221–14224. Conley, M. P.; Mougel, V.; Peryshkov, D. V.; Forrest, W. P.; Gajan, D.; Lesage, A.; Emsley, L.; Copéret, C.; Schrock, R. R. J. Am. Chem. Soc. 2013, 135, 19068–19070. Mazoyer, E.; Merle, N.; Mallmann, A. D.; Basset, J.-M.; Berrier, E.; Delevoye, L.; Paul, J.-F.; Nicholas, C. P.; Gauvin, R. M.; Taoufik, M. Chem. Commun. 2010, 46, 8944–8946.

Organometallic Chemistry on Oxide Surfaces

605

21. Merle, N.; Girard, G.; Popoff, N.; De Mallmann, A.; Bouhoute, Y.; Trébosc, J.; Berrier, E.; Paul, J.-F.; Nicholas, C. P.; Del Rosal, I.; Maron, L.; Gauvin, R. M.; Delevoye, L.; Taoufik, M. Inorg. Chem. 2013, 52, 10119–10130. 22. Merle, N.; Le Quéméner, F.; Bouhoute, Y.; Szeto, K. C.; De Mallmann, A.; Barman, S.; Samantaray, M. K.; Delevoye, L.; Gauvin, R. M.; Taoufik, M.; Basset, J.-M. J. Am. Chem. Soc. 2017, 139, 2144–2147. 23. Bouhoute, Y.; Grekov, D.; Szeto, K. C.; Merle, N.; De Mallmann, A.; Lefebvre, F.; Raffa, G.; Del Rosal, I.; Maron, L.; Gauvin, R. M.; Delevoye, L.; Taoufik, M. ACS Catal. 2016, 6, 1–18. 24. Marks, T. J. Acc. Chem. Res. 1992, 25, 57–65. 25. Wegener, S. L.; Marks, T. J.; Stair, P. C. Acc. Chem. Res. 2011, 45, 206–214. 26. Stalzer, M.; Delferro, M.; Marks, T. Catal. Lett. 2015, 145, 3–14. 27. Basset, J.-M.; Coperet, C.; Soulivong, D.; Taoufik, M.; Cazat, J. T. Acc. Chem. Res. 2009, 43, 323–334. 28. Copéret, C.; Chabanas, M.; Petroff Saint-Arroman, R.; Basset, J.-M. Angew. Chem. Int. Ed. 2003, 42, 156–181. 29. Guzman, J.; Gates, B. C. Dalton Trans. 2003, 3303–3318. https://doi.org/10.1039/b303285j. 30. Brutchey, R. L.; Tilley, T. D. Top. Organomet. Chem. 2005, 69–115. 31. Conley, M. P.; Copéret, C.; Thieuleux, C. ACS Catal. 2014, 4, 1458–1469. 32. Liang, Y.; Anwander, R. Dalton Trans. 2013, 42, 12521–12545. 33. Samantaray, M. K.; Dey, R.; Kavitake, S.; Basset, J.-M. In C-H Bond Activation and Catalytic Functionalization II; Dixneuf, P. H., Doucet, H., Eds.; Springer International Publishing: Cham, 2016; pp 155–187. https://doi.org/10.1007/3418_2015_139. 34. Copéret, C. Catal. Chem. Biol. 2018, 236–242. https://doi.org/10.1142/9789813237179_0035. 35. Applied Homogeneous Catalysis with Organometallic Compounds; pp 1069–1084. https://doi.org/10.1002/9783527651733.ch15. 36. Copéret, C.; Allouche, F.; Chang, K. W.; Conley, M.; Delley, M. F.; Fedorov, A.; Moroz, I.; Mougel, V.; Pucino, M.; Searles, K.; Yamamoto, K.; Zhizhko, P. Angew. Chem. Int. Ed. 2018, 57, 6398–6440. 37. Copéret, C.; Comas-Vives, A.; Conley, M. P.; Estes, D. P.; Fedorov, A.; Mougel, V.; Nagae, H.; Núñez-Zarur, F.; Zhizhko, P. A. Chem. Rev. 2016, 116, 323–421. 38. Samantaray, M. K.; D’Elia, V.; Pump, E.; Falivene, L.; Harb, M.; Ould Chikh, S.; Cavallo, L.; Basset, J.-M. Chem. Rev. 2020, 120, 734–813. 39. Copéret, C.; Estes, D. P.; Larmier, K.; Searles, K. Chem. Rev. 2016, 116, 8463–8505. 40. Copéret, C.; Fedorov, A.; Zhizhko, P. A. Catal. Lett. 2017, 147, 2247–2259. 41. Samantaray, M. K.; Pump, E.; Bendjeriou-Sedjerari, A.; D’Elia, V.; Pelletier, J. D. A.; Guidotti, M.; Psaro, R.; Basset, J.-M. Chem. Soc. Rev. 2018, 47, 8403–8437. 42. Grekov, D.; Vancompernolle, T.; Taoufik, M.; Delevoye, L.; Gauvin, R. M. Chem. Soc. Rev. 2018, 47, 2572–2590. 43. Copéret, C.; Liao, W.-C.; Gordon, C. P.; Ong, T.-C. J. Am. Chem. Soc. 2017, 139, 10588–10596. 44. Kobayashi, T.; Perras, F. A.; Slowing, I. I.; Sadow, A. D.; Pruski, M. ACS Catal. 2015, 5, 7055–7062. 45. Copéret, C. Nat. Energy 2019, 4, 1018–1024. 46. Goldsmith, B. R.; Peters, B.; Johnson, J. K.; Gates, B. C.; Scott, S. L. ACS Catal. 2017, 7, 7543–7557. 47. Gates, B. C.; Flytzani-Stephanopoulos, M.; Dixon, D. A.; Katz, A. Cat. Sci. Technol. 2017, 7, 4259–4275. 48. Guan, E.; Fang, C.-Y.; Yang, D.; Wang, L.; Xiao, F.-S.; Gates, B. C. Faraday Discuss. 2018, 208, 9–33. 49. Witzke, R. J.; Chapovetsky, A.; Conley, M. P.; Kaphan, D. M.; Delferro, M. ACS Catal. 2020, 11822–11840. https://doi.org/10.1021/acscatal.0c03350. 50. Pelletier, J. D. A.; Basset, J.-M. Acc. Chem. Res. 2016, 49, 664–677. 51. Copéret, C. Acc. Chem. Res. 2019, 52, 1697–1708. 52. Blanc, F.; Coperet, C.; Lesage, A.; Emsley, L. Chem. Soc. Rev. 2008, 37, 518–526. 53. Rossini, A. J.; Zagdoun, A.; Lelli, M.; Lesage, A.; Copéret, C.; Emsley, L. Acc. Chem. Res. 2013. https://doi.org/10.1021/ar300322x. 54. Rossini, A. J. J. Phys. Chem. Lett. 2018, 9, 5150–5159. 55. Bordiga, S.; Groppo, E.; Agostini, G.; van Bokhoven, J. A.; Lamberti, C. Chem. Rev. 2013, 113, 1736–1850. 56. Sautet, P.; Delbecq, F. Chem. Rev. 2010, 110, 1788–1806. 57. Chen, E. Y.-X.; Marks, T. J. Chem. Rev. 2000, 100, 1391–1434. 58. Tolman, C. A. Chem. Rev. 1977, 77, 313–348. 59. Falivene, L.; Credendino, R.; Poater, A.; Petta, A.; Serra, L.; Oliva, R.; Scarano, V.; Cavallo, L. Organometallics 2016, 35, 2286–2293. 60. Clavier, H.; Nolan, S. P. Chem. Commun. 2010, 46, 841–861. 61. Lichtenstein, L.; Büchner, C.; Yang, B.; Shaikhutdinov, S.; Heyde, M.; Sierka, M.; Włodarczyk, R.; Sauer, J.; Freund, H.-J. Angew. Chem. Int. Ed. 2012, 51, 404–407. 62. Gladden, L. F.; Carpenter, T. A.; Elliott, S. R. Philos. Mag. B 1986, 53, L81–L87. 63. Mozzi, R. L.; Warren, B. E. J. Appl. Cryst. 1969, 2, 164–172. 64. Zhuravlev, L. T. Colloids Surf. A Physicochem. Eng. Asp. 2000, 173, 1–38. 65. van der Meer, J.; Bardez-Giboire, I.; Mercier, C.; Revel, B.; Davidson, A.; Denoyel, R. J. Phys. Chem. C 2010, 114, 3507–3515. 66. Kinney, D. R.; Chuang, I. S.; Maciel, G. E. J. Am. Chem. Soc. 1993, 115, 6786–6794. 67. Maciel, G. E. J. Am. Chem. Soc. 1996, 118, 401–406. 68. Griffin, J. M.; Clark, L.; Seymour, V. R.; Aldous, D. W.; Dawson, D. M.; Iuga, D.; Morris, R. E.; Ashbrook, S. E. Chem. Sci. 2012, 3, 2293–2300. 69. Bignami, G. P. M.; Davis, Z. H.; Dawson, D. M.; Morris, S. A.; Russell, S. E.; McKay, D.; Parke, R. E.; Iuga, D.; Morris, R. E.; Ashbrook, S. E. Chem. Sci. 2018, 9, 850–859. 70. Merle, N.; Trébosc, J.; Baudouin, A.; Rosal, I. D.; Maron, L.; Szeto, K.; Genelot, M.; Mortreux, A.; Taoufik, M.; Delevoye, L.; Gauvin, R. M. J. Am. Chem. Soc. 2012, 134, 9263–9275. 71. Poater, A.; Solans-Monfort, X.; Clot, E.; Copéret, C.; Eisenstein, O. J. Am. Chem. Soc. 2007, 129, 8207–8216. 72. Solans-Monfort, X.; Clot, E.; Copéret, C.; Eisenstein, O. J. Am. Chem. Soc. 2005, 127, 14015–14025. 73. Estes, D. P.; Gordon, C. P.; Fedorov, A.; Liao, W. C.; Ehrhorn, H.; Bittner, C.; Zier, M. L.; Bockfeld, D.; Chan, K. W.; Eisenstein, O.; Raynaud, C.; Tamm, M.; Coperet, C. J. Am. Chem. Soc. 2017, 139, 17597–17607. 74. Halbert, S.; Coperet, C.; Raynaud, C.; Eisenstein, O. J. Am. Chem. Soc. 2016, 138, 2261–2272. 75. Del Rosal, I.; Gerber, I. C.; Poteau, R.; Maron, L. J. Phys. Chem. A 2010, 114, 6322–6330. 76. Culver, D. B.; Venkatesh, A.; Huynh, W.; Rossini, A. J.; Conley, M. P. Chem. Sci. 2020, 11, 1510–1517. 77. Culver, D. B.; Huynh, W.; Tafazolian, H.; Conley, M. P. Organometallics 2020, 39, 1112–1122. 78. Solans-Monfort, X.; Filhol, J.-S.; Copéret, C.; Eisenstein, O. New J. Chem. 2006, 30, 842–850. 79. Comas-Vives, A. Phys. Chem. Chem. Phys. 2016, 18, 7475–7482. 80. Ewing, C. S.; Bhavsar, S.; Veser, G.; McCarthy, J. J.; Johnson, J. K. Langmuir 2014, 30, 5133–5141. 81. Zecchina, A.; Scarano, D.; Bordiga, S.; Spoto, G.; Lamberti, C. Adv. Catal. 2001, 46, 265–397. 82. Busca, G. Catal. Today 2014, 226, 2–13. 83. Wischert, R.; Laurent, P.; Copéret, C.; Delbecq, F.; Sautet, P. J. Am. Chem. Soc. 2012, 134, 14430–14449. 84. Martin, D.; Duprez, D. J. Phys. Chem. 1996, 100, 9429–9438. 85. Taoufik, M.; Szeto, K. C.; Merle, N.; Rosal, I. D.; Maron, L.; Trébosc, J.; Tricot, G.; Gauvin, R. M.; Delevoye, L. Chem. A Eur. J. 2014, 20, 4038–4046. 86. Wischert, R.; Copéret, C.; Delbecq, F.; Sautet, P. Angew. Chem. Int. Ed. 2011, 50, 3202–3205.

606

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. 131. 132. 133. 134. 135. 136. 137. 138. 139. 140. 141. 142. 143. 144. 145. 146. 147. 148. 149. 150. 151. 152. 153. 154.

Organometallic Chemistry on Oxide Surfaces

Wischert, R.; Coperet, C.; Delbecq, F.; Sautet, P. Chem. Commun. 2011, 47, 4890–4892. Digne, M.; Sautet, P.; Raybaud, P.; Euzen, P.; Toulhoat, H. J. Catal. 2002, 211, 1–5. Comas-Vives, A.; Schwarzwälder, M.; Copéret, C.; Sautet, P. J. Phys. Chem. C 2015, 119, 7156–7163. Comas-Vives, A.; Valla, M.; Copéret, C.; Sautet, P. ACS Cent. Sci. 2015. https://doi.org/10.1021/acscentsci.5b00226. Hensen, E. J. M.; Poduval, D. G.; Magusin, P. C. M. M.; Coumans, A. E.; van Veen, J. A. R. J. Catal. 2010, 269, 201–218. Corma, A. Chem. Rev. 1995, 95, 559–614. Miller, J. B.; Ko, E. I. Catal. Today 1997, 35, 269–292. La Parola, V.; Deganello, G.; Scirè, S.; Venezia, A. M. J. Solid State Chem. 2003, 174, 482–488. Omegna, A.; van Bokhoven, J. A.; Prins, R. J. Phys. Chem. B 2003, 107, 8854–8860. Caillot, M.; Chaumonnot, A.; Digne, M.; Poleunis, C.; Debecker, D. P.; van Bokhoven, J. A. Microp. Mesop. Mater. 2014, 185, 179–189. Caillot, M.; Chaumonnot, A.; Digne, M.; van Bokhoven, J. A. J. Catal. 2014, 316, 47–56. Mouat, A. R.; Kobayashi, T.; Pruski, M.; Marks, T. J.; Stair, P. C. J. Phys. Chem. C 2017, 121, 6060–6064. Mouat, A. R.; George, C.; Kobayashi, T.; Pruski, M.; van Duyne, R. P.; Marks, T. J.; Stair, P. C. Angew. Chem. Int. Ed. 2015, 54, 13346–13351. Crépeau, G.; Montouillout, V.; Vimont, A.; Mariey, L.; Cseri, T.; Maugé, F. J. Phys. Chem. B 2006, 110, 15172–15185. Trombetta, M.; Busca, G.; Rossini, S.; Piccoli, V.; Cornaro, U.; Guercio, A.; Catani, R.; Willey, R. J. J. Catal. 1998, 179, 581–596. Hensen, E. J. M.; Poduval, D. G.; Degirmenci, V.; Ligthart, D. A. J. M.; Chen, W.; Maugé, F.; Rigutto, M. S.; van Veen, J. A. R. J. Phys. Chem. C 2012, 116, 21416–21429. Busca, G. Phys. Chem. Chem. Phys. 1999, 1, 723–736. Daniell, W.; Schubert, U.; Glöckler, R.; Meyer, A.; Noweck, K.; Knözinger, H. Appl. Catal. A. Gen. 2000, 196, 247–260. Finocchio, E.; Busca, G.; Rossini, S.; Cornaro, U.; Piccoli, V.; Miglio, R. Catal. Today 1997, 33, 335–352. Valla, M.; Rossini, A. J.; Caillot, M.; Chizallet, C.; Raybaud, P.; Digne, M.; Chaumonnot, A.; Lesage, A.; Emsley, L.; van Bokhoven, J. A.; Copéret, C. J. Am. Chem. Soc. 2015, 137, 10710–10719. Hino, M.; Kobayashi, S.; Arata, K. J. Am. Chem. Soc. 1979, 101, 6439–6441. Hino, M.; Arata, K. Chem. Commun. 1980, 851–852. https://doi.org/10.1039/c39800000851. Chen, W.; Yi, X.; Huang, L.; Liu, W.; Li, G.; Acharya, D.; Sun, X.; Zheng, A. Cat. Sci. Technol. 2019, 9, 5045–5057. Tabora, J. E.; Davis, R. J. J. Am. Chem. Soc. 1996, 118, 12240–12241. Li, X.; Nagaoka, K.; Simon, L. J.; Olindo, R.; Lercher, J. A.; Hofmann, A.; Sauer, J. J. Am. Chem. Soc. 2005, 127, 16159–16166. Rodriguez, J.; Culver, D. B.; Conley, M. P. J. Am. Chem. Soc. 2019, 141, 1484–1488. Farneth, W. E.; Gorte, R. J. Chem. Rev. 1995, 95, 615–635. Gorte, R. J. Catal. Lett. 1999, 62, 1–13. Haase, F.; Sauer, J. J. Am. Chem. Soc. 1998, 120, 13503–13512. Viggiano, A. A.; Henchman, M. J.; Dale, F.; Deakyne, C. A.; Paulson, J. F. J. Am. Chem. Soc. 1992, 114, 4299–4306. Sauer, J.; Ahlrichs, R. J. Chem. Phys. 1990, 93, 2575–2583. Williams, L. A.; Guo, N.; Motta, A.; Delferro, M.; Fragala, I. L.; Miller, J. T.; Marks, T. J. Proc. Natl. Acad. Sci. U. S. A. 2013, 110, 413–418. Grinenval, E.; Bayard, F.; Basset, J.-M.; Lefebvre, F. Inorg. Chem. 2014, 53, 2022–2029. Grinenval, E.; Rozanska, X.; Baudouin, A.; Berrier, E.; Delbecq, F.; Sautet, P.; Basset, J.-M.; Lefebvre, F. J. Phys. Chem. C 2010, 114, 19024–19034. Palermo, A. P.; Schöttle, C.; Zhang, S.; Grosso-Giordano, N. A.; Okrut, A.; Dixon, D. A.; Frei, H.; Gates, B. C.; Katz, A. Inorg. Chem. 2019, 58, 14338–14348. Palermo, A.; Solovyov, A.; Ertler, D.; Okrut, A.; Gates, B. C.; Katz, A. Chem. Sci. 2017, 8, 4951–4960. Palermo, A. P.; Zhang, S.; Hwang, S.-J.; Dixon, D. A.; Gates, B. C.; Katz, A. Dalton Trans. 2018, 47, 13550–13558. Schöttle, C.; Guan, E.; Okrut, A.; Grosso-Giordano, N. A.; Palermo, A.; Solovyov, A.; Gates, B. C.; Katz, A. J. Am. Chem. Soc. 2019, 141, 4010–4015. Kharel, S.; Cluff, K. J.; Bhuvanesh, N.; Gladysz, J. A.; Blümel, J. Chem. Asian J. 2019, 14, 2704–2711. Dufaud, V.; Niccolai, G. P.; Thivolle-Cazat, J.; Basset, J.-M. J. Am. Chem. Soc. 1995, 117, 4288–4294. LaPointe, R. E.; Wolczanski, P. T.; Van Duyne, G. D. Organometallics 1985, 4, 1810–1818. Quadrelli, E. A.; Basset, J. M. Coord. Chem. Rev. 2010, 254, 707–728. Gajan, D.; Rendón, N.; Wampler, K. M.; Jean-Marie, B.; Copéret, C.; Lesage, A.; Emsley, L.; Schrock, R. R. Dalton Trans. 2010, 39, 8547–8551. Conley, M. P.; Lapadula, G.; Sanders, K.; Gajan, D.; Lesage, A.; del Rosal, I.; Maron, L.; Lukens, W. W.; Copéret, C.; Andersen, R. A. J. Am. Chem. Soc. 2016, 138, 3831–3843. Le Roux, E.; Chabanas, M.; Baudouin, A.; de Mallmann, A.; Copéret, C.; Quadrelli, E. A.; Thivolle-Cazat, J.; Basset, J.-M.; Lukens, W.; Lesage, A.; Emsley, L.; Sunley, G. J. J. Am. Chem. Soc. 2004, 126, 13391–13399. Lesage, A.; Emsley, L.; Chabanas, M.; Copéret, C.; Basset, J.-M. Angew. Chem. Int. Ed. 2002, 41, 4535–4538. Rossini, A. J.; Zagdoun, A.; Lelli, M.; Lesage, A.; Copéret, C.; Emsley, L. Acc. Chem. Res. 2013, 46, 1942–1951. Pump, E.; Viger-Gravel, J.; Abou-Hamad, E.; Samantaray, M. K.; Hamzaoui, B.; Gurinov, A.; Anjum, D. H.; Gajan, D.; Lesage, A.; Bendjeriou-Sedjerari, A.; Emsley, L.; Basset, J.-M. Chem. Sci. 2017, 8, 284–290. Ong, T.-C.; Liao, W.-C.; Mougel, V.; Gajan, D.; Lesage, A.; Emsley, L.; Copéret, C. Angew. Chem. Int. Ed. 2016, 55, 4743–4747. Perras, F. A.; Boteju, K. C.; Slowing, I. I.; Sadow, A. D.; Pruski, M. Chem. Commun. 2018, 54, 3472–3475. Autschbach, J.; Zheng, S.; Schurko, R. W. Concepts Magn. Reson. A 2010, 36A, 84–126. Vancompernolle, T.; Trivelli, X.; Delevoye, L.; Pourpoint, F.; Gauvin, R. M. Dalton Trans. 2017, 46, 13176–13179. Wolczanski, P. T.; Chirik, P. J. ACS Catal. 2015, 5, 1747–1757. Culver, D. B.; Huynh, W.; Tafazolian, H.; Ong, T.-C.; Conley, M. P. Angew. Chem. Int. Ed. 2018, 57, 9520–9523. Huynh, W.; Culver, D. B.; Tafazolian, H.; Conley, M. P. Dalton Trans. 2018, 47, 13063–13071. Rossini, A. J.; Schurko, R. W. J. Am. Chem. Soc. 2006, 128, 10391–10402. Giovine, R.; Volkringer, C.; Ashbrook, S. E.; Trébosc, J.; McKay, D.; Loiseau, T.; Amoureux, J.-P.; Lafon, O.; Pourpoint, F. Chem. A Eur. J. 2017, 23, 9525–9534. Avent, A. G.; Caro, C. F.; Hitchcock, P. B.; Lappert, M. F.; Li, Z.; Wei, X.-H. Dalton Trans. 2004, 1567–1577. https://doi.org/10.1039/b316695n. Hitchcock, P. B.; Lappert, M. F.; Smith, R. G.; Bartlett, R. A.; Power, P. P. Chem. Commun. 1988, 1007–1009. https://doi.org/10.1039/c39880001007. Eedugurala, N.; Wang, Z.; Yan, K.; Boteju, K. C.; Chaudhary, U.; Kobayashi, T.; Ellern, A.; Slowing, I. I.; Pruski, M.; Sadow, A. D. Organometallics 2017, 36, 1142–1153. Pindwal, A.; Yan, K.; Patnaik, S.; Schmidt, B. M.; Ellern, A.; Slowing, I. I.; Bae, C.; Sadow, A. D. J. Am. Chem. Soc. 2017, 139, 16862–16874. Wang, Z.; Patnaik, S.; Eedugurala, N.; Manzano, J. S.; Slowing, I. I.; Kobayashi, T.; Sadow, A. D.; Pruski, M. J. Am. Chem. Soc. 2020, 142, 2935–2947. Green, M. L. H.; Parkin, G. J. Chem. Educ. 2014, 91, 807–816. Vancompernolle, T.; Merle, N.; Capet, F.; Del Rosal, I.; Laurent, M.; Delevoye, L.; Pourpoint, F.; Gauvin, R. M. Dalton Trans. 2019, 48, 5243–5252. Gajan, D.; Coperet, C. New J. Chem. 2011, 35, 2403–2408. LaPointe, A. M.; Schrock, R. R. Organometallics 1993, 12, 3379–3381. Berthoud, R.; Rendon, N.; Blanc, F.; Solans-Monfort, X.; Coperet, C.; Eisenstein, O. Dalton Trans. 2009, 5879–5886. Rhers, B.; Quadrelli, E. A.; Baudouin, A.; Taoufik, M.; Copéret, C.; Lefebvre, F.; Basset, J.-M.; Fenet, B.; Sinha, A.; Schrock, R. R. J. Organomet. Chem. 2006, 691, 5448–5455.

Organometallic Chemistry on Oxide Surfaces

607

155. Lefort, L.; Chabanas, M.; Maury, O.; Meunier, D.; Copéret, C.; Thivolle-Cazat, J.; Basset, J.-M. J. Organomet. Chem. 2000, 593-594, 96–100. 156. Quignard, F.; Lecuyer, C.; Bougault, C.; Lefebvre, F.; Choplin, A.; Olivier, D.; Basset, J. M. Inorg. Chem. 1992, 31, 928–930. 157. Langeslay, R. R.; Sohn, H.; Hu, B.; Mohar, J. S.; Ferrandon, M.; Liu, C.; Kim, H.; Jeremy Kropf, A.; Yang, C.; Niklas, J.; Poluektov, O. G.; Ercan Alp, E.; Ignacio-de Leon, P.; Sattelberger, A. P.; Hock, A. S.; Delferro, M. Dalton Trans. 2018, 47, 10842–10846. 158. Merle, N.; Le Quéméner, F.; Barman, S.; Samantaray, M. K.; Szeto, K. C.; De Mallmann, A.; Taoufik, M.; Basset, J.-M. Chem. Commun. 2017, 53, 11338–11341. 159. Grekov, D.; Bouhoute, Y.; Szeto, K. C.; Merle, N.; De Mallmann, A.; Lefebvre, F.; Lucas, C.; Del Rosal, I.; Maron, L.; Gauvin, R. M.; Delevoye, L.; Taoufik, M. Organometallics 2016, 35, 2188–2196. 160. Rouge, P.; Szeto, K. C.; Bouhoute, Y.; Merle, N.; De Mallmann, A.; Delevoye, L.; Gauvin, R. M.; Taoufik, M. Organometallics 2020, 39, 1105–1111. 161. Copéret, C.; Basset, J.-M. Adv. Synth. Catal. 2007, 349, 78–92. 162. Popoff, N.; Mazoyer, E.; Pelletier, J.; Gauvin, R. M.; Taoufik, M. Chem. Soc. Rev. 2013, 42, 9035–9054. 163. Gordon, C. P.; Yamamoto, K.; Liao, W. C.; Allouche, F.; Andersen, R. A.; Coperet, C.; Raynaud, C.; Eisenstein, O. ACS Cent. Sci. 2017, 3, 759–768. 164. Mougel, V.; Santiago, C. B.; Zhizhko, P. A.; Bess, E. N.; Varga, J.; Frater, G.; Sigman, M. S.; Copéret, C. J. Am. Chem. Soc. 2015, 137, 6699–6704. 165. De Jesus Silva, J.; Ferreira, M. A. B.; Fedorov, A.; Sigman, M. S.; Copéret, C. Chem. Sci. 2020, 11, 6717–6723. 166. Hoveyda, A. H.; Liu, Z.; Qin, C.; Koengeter, T.; Mu, Y. Angew. Chem. Int. Ed. 59, 2020, 22324–22348. 167. Flook, M. M.; Jiang, A. J.; Schrock, R. R.; Müller, P.; Hoveyda, A. H. J. Am. Chem. Soc. 2009, 131, 7962–7963. 168. Townsend, E. M.; Schrock, R. R.; Hoveyda, A. H. J. Am. Chem. Soc. 2012, 134, 11334–11337. 169. Lopez, L. P. H.; Schrock, R. R.; Muller, P. Organometallics 2006, 25, 1978–1986. 170. Schrock, R. R. Chem. Rev. 2009, 109, 3211–3226. 171. Peryshkov, D. V.; Forrest, W. P.; Schrock, R. R.; Smith, S. J.; Müller, P. Organometallics 2013, 32, 5256–5259. 172. Lassalle, S.; Jabbour, R.; Schiltz, P.; Berruyer, P.; Todorova, T. K.; Veyre, L.; Gajan, D.; Lesage, A.; Thieuleux, C.; Camp, C. J. Am. Chem. Soc. 2019, 141, 19321–19335. 173. Nicholas, C. P.; Ahn, H.; Marks, T. J. J. Am. Chem. Soc. 2003, 125, 4325–4331. 174. Fandos, R.; Hernández, C.; Otero, A.; Rodríguez, A.; Ruiz, M. J.; Terreros, P. J. Organomet. Chem. 2000, 606, 156–162. 175. Yang, X.; Stern, C. L.; Marks, T. J. J. Am. Chem. Soc. 1994, 116, 10015–10031. 176. Ittel, S. D.; Johnson, L. K.; Brookhart, M. Chem. Rev. 2000, 100, 1169–1204. 177. Williams, L. A.; Marks, T. J. ACS Catal. 2011, 1, 238–245. 178. Stalzer, M. M.; Nicholas, C. P.; Bhattacharyya, A.; Motta, A.; Delferro, M.; Marks, T. J. Angew. Chem. Int. Ed. 2016, 55, 5263–5267. 179. Nicholas, C. P.; Marks, T. J. Langmuir 2004, 20, 9456–9462. 180. Nicholas, C. P.; Marks, T. J. Nano Lett. 2004, 4, 1557–1559. 181. Ahn, H.; Nicholas, C. P.; Marks, T. J. Organometallics 2002, 21, 1788–1806. 182. Gao, Y.; Mouat, A. R.; Motta, A.; Macchioni, A.; Zuccaccia, C.; Delferro, M.; Marks, T. J. ACS Catal. 2015, 5, 5272–5282. 183. Zhang, J.; Motta, A.; Gao, Y.; Stalzer, M. M.; Delferro, M.; Liu, B.; Lohr, T. L.; Marks, T. J. ACS Catal. 2018, 8, 4893–4901. 184. Gu, W.; Stalzer, M. M.; Nicholas, C. P.; Bhattacharyya, A.; Motta, A.; Gallagher, J. R.; Zhang, G.; Miller, J. T.; Kobayashi, T.; Pruski, M.; Delferro, M.; Marks, T. J. J. Am. Chem. Soc. 2015, 137, 6770–6780. 185. Culver, D. B.; Tafazolian, H.; Conley, M. P. Organometallics 2018, 37, 1001–1006. 186. Tafazolian, H.; Culver, D. B.; Conley, M. P. Organometallics 2017, 36, 2385–2388. 187. Kaphan, D. M.; Klet, R. C.; Perras, F. A.; Pruski, M.; Yang, C.; Kropf, A. J.; Delferro, M. ACS Catal. 2018, 8, 5363–5373. 188. Syed, Z. H.; Kaphan, D. M.; Perras, F. A.; Pruski, M.; Ferrandon, M. S.; Wegener, E. C.; Celik, G.; Wen, J.; Liu, C.; Dogan, F.; Goldberg, K. I.; Delferro, M. J. Am. Chem. Soc. 2019, 141, 6325–6337. 189. Culver, D. B.; Conley, M. P. Angew. Chem. Int. Ed. 2018, 57, 14902–14905. 190. Lambert, J. B.; Zhang, S. Science 1994, 263, 984–985. 191. Reed, C. A.; Xie, Z. Science 1994, 263, 985–986. 192. Pauling, L. Science 1994, 263, 983–984. 193. Olah, G. A.; Rasul, G.; Li, X.-y.; Buchholz, H. A.; Sandford, G.; Prakash, G. K. S. Science 1994, 263, 983–984. 194. Kim, K.-C.; Reed, C. A.; Elliott, D. W.; Mueller, L. J.; Tham, F.; Lin, L.; Lambert, J. B. Science 2002, 297, 825–827. 195. Reed, C. A. Acc. Chem. Res. 1998, 31, 325–332. 196. Reed, C. A. Acc. Chem. Res. 2010, 43, 121–128. 197. Huynh, W.; Conley, M. P. Dalton Trans. 2020, 49, 16453–16463. 198. Klare, H. F. T. ACS Catal. 2017, 7, 6999–7002. 199. Deschner, T.; Liang, Y.; Anwander, R. J. Phys. Chem. C 2010, 114, 22603–22609. 200. Zapilko, C.; Widenmeyer, M.; Nagl, I.; Estler, F.; Anwander, R.; Raudaschl-Sieber, G.; Groeger, O.; Engelhardt, G. J. Am. Chem. Soc. 2006, 128, 16266–16276. 201. Bluemel, J. J. Am. Chem. Soc. 1995, 117, 2112–2113. 202. Ahrens, M.; Scholz, G.; Braun, T.; Kemnitz, E. Angew. Chem. Int. Ed. 2013, 52, 5328–5332. 203. Scott, V. J.; Celenligil-Cetin, R.; Ozerov, O. V. J. Am. Chem. Soc. 2005, 127, 2852–2853. 204. Douvris, C.; Ozerov, O. V. Science 2008, 321, 1188–1190. 205. Shao, B.; Bagdasarian, A. L.; Popov, S.; Nelson, H. M. Science 2017, 355, 1403–1407. 206. Klet, R. C.; Kaphan, D. M.; Liu, C.; Yang, C.; Kropf, A. J.; Perras, F. A.; Pruski, M.; Hock, A. S.; Delferro, M. J. Am. Chem. Soc. 2018, 140, 6308–6316. 207. Stoyanov, E. S.; Kim, K.-C.; Reed, C. A. J. Am. Chem. Soc. 2006, 128, 8500–8508. 208. Yamamoto, H.; Futatsugi, K. Angew. Chem. Int. Ed. 2005, 44, 1924–1942. 209. Müller, L. O.; Himmel, D.; Stauffer, J.; Steinfeld, G.; Slattery, J.; Santiso-Quiñones, G.; Brecht, V.; Krossing, I. Angew. Chem. Int. Ed. 2008, 47, 7659–7663. 210. Bendjeriou-Sedjerari, A.; Azzi, J. M.; Abou-Hamad, E.; Anjum, D. H.; Pasha, F. A.; Huang, K.-W.; Emsley, L.; Basset, J.-M. J. Am. Chem. Soc. 2013, 135, 17943–17951. 211. Morrow, B. A.; Cody, I. A.; Lee, L. S. M. J. Phys. Chem. 1976, 80, 2761–2767. 212. Conley, M. P.; Rossini, A. J.; Comas-Vives, A.; Valla, M.; Casano, G.; Ouari, O.; Tordo, P.; Lesage, A.; Emsley, L.; Copéret, C. Phys. Chem. Chem. Phys. 2014, 16, 17822–17827. 213. Kermagoret, A.; Kerber, R. N.; Conley, M. P.; Callens, E.; Florian, P.; Massiot, D.; Coperet, C.; Delbecq, F.; Rozanska, X.; Sautet, P. Dalton Trans. 2013, 42, 12681–12687. 214. Kerber, R. N.; Kermagoret, A.; Callens, E.; Florian, P.; Massiot, D.; Lesage, A.; Copéret, C.; Delbecq, F.; Rozanska, X.; Sautet, P. J. Am. Chem. Soc. 2012, 134, 6767–6775. 215. Kermagoret, A.; Kerber, R. N.; Conley, M. P.; Callens, E.; Florian, P.; Massiot, D.; Delbecq, F.; Rozanska, X.; Copéret, C.; Sautet, P. J. Catal. 2014, 313, 46–54. 216. Moroz, I. B.; Florian, P.; Viger-Gravel, J.; Gordon, C. P.; Lesage, A.; Copéret, C. Angew. Chem. Int. Ed. 2020, 59, 16167–16172. 217. Bashir, M. A.; Vancompernolle, T.; Gauvin, R. M.; Delevoye, L.; Merle, N.; Monteil, V.; Taoufik, M.; McKenna, T. F. L.; Boisson, C. Cat. Sci. Technol. 2016, 6, 2962–2974. 218. Sauter, D. W.; Popoff, N.; Bashir, M. A.; Szeto, K. C.; Gauvin, R. M.; Delevoye, L.; Taoufik, M.; Boisson, C. Chem. Commun. 2016, 52, 4776–4779. 219. Pelletier, J.; Espinas, J.; Vu, N.; Norsic, S.; Baudouin, A.; Delevoye, L.; Trébosc, J.; Le Roux, E.; Santini, C.; Basset, J.-M.; Gauvin, R. M.; Taoufik, M. Chem. Commun. 2011, 47, 2979–2981. 220. Sauter, D. W.; Chiari, V.; Aykac, N.; Bouaouli, S.; Perrin, L.; Delevoye, L.; Gauvin, R. M.; Szeto, K. C.; Boisson, C.; Taoufik, M. Dalton Trans. 2017, 46, 11547–11551. 221. Fleischman, S. D.; Scott, S. L. J. Am. Chem. Soc. 2011, 133, 4847–4855.

608

222. 223. 224. 225. 226. 227. 228. 229. 230. 231. 232. 233. 234. 235. 236. 237. 238. 239. 240. 241. 242. 243. 244. 245. 246. 247. 248. 249. 250. 251. 252. 253. 254. 255. 256. 257. 258. 259. 260. 261. 262. 263. 264. 265. 266. 267. 268.

Organometallic Chemistry on Oxide Surfaces

Taha, Z. A.; Deguns, E. W.; Chattopadhyay, S.; Scott, S. L. Organometallics 2006, 25, 1891–1899. Szeto, K. C.; Jones, Z. R.; Merle, N.; Rios, C.; Gallo, A.; Le Quemener, F.; Delevoye, L.; Gauvin, R. M.; Scott, S. L.; Taoufik, M. ACS Catal. 2018, 8, 7566–7577. Demmelmaier, C. A.; White, R. E.; van Bokhoven, J. A.; Scott, S. L. J. Phys. Chem. C 2008, 116, 6439–6449. McDaniel, M. P. J. Catal. 1982, 76, 17–28. Scott, S. L.; Basset, J.-M. J. Am. Chem. Soc. 1994, 116, 12069–12070. Mania, P.; Verel, R.; Jenny, F.; Hammond, C.; Hermans, I. Chem. A Eur. J. 2013, 19, 9849–9858. Vancompernolle, T.; Valente, A.; Chenal, T.; Zinck, P.; Del Rosal, I.; Maron, L.; Taoufik, M.; Harder, S.; Gauvin, R. M. Organometallics 2017, 36, 3912–3920. Vidal, V.; Théolier, A.; Thivolle-Cazat, J.; Basset, J.-M. Science 1997, 276, 99–102. Corker, J.; Lefebvre, F.; Lécuyer, C.; Dufaud, V.; Quignard, F.; Choplin, A.; Evans, J.; Basset, J.-M. Science 1996, 271, 966–969. Avenier, P.; Taoufik, M.; Lesage, A.; Solans-Monfort, X.; Baudouin, A.; de Mallmann, A.; Veyre, L.; Basset, J.-M.; Eisenstein, O.; Emsley, L.; Quadrelli, E. A. Science 2007, 317, 1056–1060. Chen, Y.; Abou-hamad, E.; Hamieh, A.; Hamzaoui, B.; Emsley, L.; Basset, J.-M. J. Am. Chem. Soc. 2015, 137, 588–591. Samantaray, M. K.; Callens, E.; Abou-Hamad, E.; Rossini, A. J.; Widdifield, C. M.; Dey, R.; Emsley, L.; Basset, J.-M. J. Am. Chem. Soc. 2014, 136, 1054–1061. Fujdala, K. L.; Tilley, T. D. J. Catal. 2003, 216, 265–275. Terry, K. W.; Lugmair, C. G.; Tilley, T. D. J. Am. Chem. Soc. 1997, 119, 9745–9756. Terry, K. W.; Ganzel, P. K.; Tilley, T. D. Chem. Mater. 1992, 4, 1290–1295. Coles, M. P.; Lugmair, C. G.; Terry, K. W.; Tilley, T. D. Chem. Mater. 2000, 12, 122–131. Brutchey, R. L.; Lugmair, C. G.; Schebaum, L. O.; Tilley, T. D. J. Catal. 2005, 229, 72–81. Brutchey, R. L.; Mork, B. V.; Sirbuly, D. J.; Yang, P. D.; Tilley, T. D. J. Mol. Catal. A Chem. 2005, 238, 1–12. Nozaki, C.; Lugmair, C. G.; Bell, A. T.; Tilley, T. D. J. Am. Chem. Soc. 2002, 124, 13194–13203. Ruddy, D. A.; Tilley, T. D. J. Am. Chem. Soc. 2008, 130, 11088–11096. Lapadula, G.; Bourdolle, A.; Allouche, F.; Conley, M. P.; del Rosal, I.; Maron, L.; Lukens, W. W.; Guyot, Y.; Andraud, C.; Brasselet, S.; Copéret, C.; Maury, O.; Andersen, R. A. Chem. Mater. 2013, 26, 1062–1073. Delley, M. F.; Lapadula, G.; Núñez-Zarur, F.; Comas-Vives, A.; Kalendra, V.; Jeschke, G.; Baabe, D.; Walter, M. D.; Rossini, A. J.; Lesage, A.; Emsley, L.; Maury, O.; Copéret, C. J. Am. Chem. Soc. 2017, 139, 8855–8867. Ballard, D. G. H. J. Polymer Sci. Polymer Chem. Ed. 1975, 13, 2191–2212. Ballard, D. G. H.; Jones, E.; Wyatt, R. J.; Murray, R. T.; Robinson, P. A. Polymer 1974, 15, 169–174. Ballard, D. G. H.; Eley, H. P. D. D.; Paul, B. W. Advances in Catalysis; Academic Press, 1973; vol. 23 pp 263–325. Yermakov, Y. I. In Studies in Surface Science and Catalysis; Seivama, T., Tanabe, K., Eds.; Elsevier, 1981; vol. 7 pp 57–76. Part A. Zakharov, V. A.; Dudchenko, V. K.; Paukshtis, E. A.; Karakchiev, L. G.; Yermakov, Y. I. J. Mol. Catal. 1977, 2, 421–435. Burwell, R. L., Jr. J. Catal. 1984, 86, 301–314. He, M. Y.; Xiong, G.; Toscano, P. J.; Burwell, R. L.; Marks, T. J. J. Am. Chem. Soc. 1985, 107, 641–652. Toscano, P. J.; Marks, T. J. J. Am. Chem. Soc. 1985, 107, 653–659. Lin, Z.; Le Marechal, J. F.; Sabat, M.; Marks, T. J. J. Am. Chem. Soc. 1987, 109, 4127–4129. Jordan, R. F.; Bajgur, C. S.; Willett, R.; Scott, B. J. Am. Chem. Soc. 1986, 108, 7410–7411. Jordan, R. F.; Dasher, W. E.; Echols, S. F. J. Am. Chem. Soc. 1986, 108, 1718–1719. Hlatky, G. G.; Turner, H. W.; Eckman, R. R. J. Am. Chem. Soc. 1989, 111, 2728–2729. Jezequel, M.; Dufaud, V.; Ruiz-Garcia, M. J.; Carrillo-Hermosilla, F.; Neugebauer, U.; Niccolai, G. P.; Lefebvre, F.; Bayard, F.; Corker, J.; Fiddy, S.; Evans, J.; Broyer, J.-P.; Malinge, J.; Basset, J.-M. J. Am. Chem. Soc. 2001, 123, 3520–3540. Joubert, J.; Delbecq, F.; Sautet, P.; Le Roux, E.; Taoufik, M.; Thieuleux, C.; Blanc, F.; Coperet, C.; Thivolle-Cazat, J.; Basset, J. M. J. Am. Chem. Soc. 2006, 128, 9157–9169. Delgado, M.; Santini, C. C.; Delbecq, F.; Wischert, R.; Le Guennic, B.; Tosin, G.; Spitz, R.; Basset, J.-M.; Sautet, P. J. Phys. Chem. C 2010, 114, 18516–18528. Delgado, M.; Delbecq, F.; Santini, C. C.; Lefebvre, F.; Norsic, S.; Putaj, P.; Sautet, P.; Basset, J.-M. J. Phys. Chem. C 2012, 116, 834–843. Taoufik, M.; Le Roux, E.; Thivolle-Cazat, J.; Copéret, C.; Basset, J.-M.; Maunders, B.; Sunley, G. J. Top. Catal. 2006, 40, 65–70. Ahn, H.; Marks, T. J. J. Am. Chem. Soc. 2002, 124, 7103–7110. Le Roux, E.; Taoufik, M.; Coperet, C.; de Mallmann, A.; Thivolle-Cazat, J.; Basset, J. M.; Maunders, B. M.; Sunley, G. J. Angew. Chem. Int. Ed. 2005, 44, 6755–6758. Coperet, C. Chem. Rev. 2010, 110, 656–680. Salameh, A.; Joubert, J.; Baudouin, A.; Lukens, W.; Delbecq, F.; Sautet, P.; Basset, J. M.; Copéret, C. Angew. Chem. Int. Ed. 2007, 46, 3870–3873. Valla, M.; Wischert, R.; Comas-Vives, A.; Conley, M. P.; Verel, R.; Copéret, C.; Sautet, P. J. Am. Chem. Soc. 2016, 138, 6774–6785. Valla, M.; Conley, M. P.; Copéret, C. Cat. Sci. Technol. 2015, 5, 1438–1442. Zhang, F.; Szeto, K. C.; Taoufik, M.; Delevoye, L.; Gauvin, R. M.; Scott, S. L. J. Am. Chem. Soc. 2018, 140, 13854–13868. Gallo, A.; Fong, A.; Szeto, K. C.; Rieb, J.; Delevoye, L.; Gauvin, R. M.; Taoufik, M.; Peters, B.; Scott, S. L. J. Am. Chem. Soc. 2016, 138, 12935–12947.

1.20

Separation Strategies in Organometallic Catalysis

Fernanda G Mendonça and R Tom Baker, Department of Chemistry and Biomolecular Sciences and Centre for Catalysis Research and Innovation, University of Ottawa, Ottawa, ON, Canada © 2022 Elsevier Ltd. All rights reserved.

1.20.1 1.20.2 1.20.2.1 1.20.2.2 1.20.2.3 1.20.2.4 1.20.2.5 1.20.2.6 1.20.3 1.20.3.1 1.20.3.2 1.20.3.3 1.20.3.4 1.20.3.5 1.20.4 References

1.20.1

Introduction Molecular catalyst immobilization on supports Metal oxide supports Carbon supports Metal-organic frameworks Polymer supports Supported ionic liquid phase Concluding remarks Multiphase strategies for catalyst separation Immobilization in polar and non-polar phases Immobilization in ionic liquids Fluorous biphasic catalysis Thermomorphic multicomponent solvent systems Concluding remarks Conclusion and outlook

609 609 610 617 619 621 624 626 626 626 628 630 631 632 632 632

Introduction

Over the last four decades, organometallic catalysis has generated a host of scientific and industrially significant advances that include Nobel Prizes for asymmetric catalysis, alkene metathesis and C-C cross-coupling and a revolution in controlled hydrocarbon polymer synthesis. Clever ligand designs, new reaction mechanisms bolstered by creative computational chemistry, and innovative high-throughput experimentation have all pushed the activity and selectivity of molecular catalysts, increasingly including non-precious metals, to new levels. Nonetheless, effective general solutions are not yet available to address the “elephant in the room,” namely efficient separation of the catalyst from the reaction products for reuse/recycle, yielding metal-free products. This topic has been elegantly summarized in several monographs dedicated to applied1–3 and multiphase4 homogeneous catalysis, catalysis in ionic liquids,5 and several recent reviews on supported molecular catalysts.6–10 In 2007, Comprehensive Organometallic Chemistry-III included special topics on chemistry in aqueous and biphasic media and in ionic liquids.11 Catalyst separation strategies can be conveniently grouped into two categories: immobilization on supports and multiphase catalysis (Fig. 1). This article covers the landscape of recent approaches to molecular catalyst separation with a focus on the last 12 years.

1.20.2

Molecular catalyst immobilization on supports

Immobilization of molecular organometallic catalysts on supports can potentially combine the advantages of homogeneous and heterogeneous catalysis, especially concerning catalyst separation and recycling. Supported molecular catalysts can also be applied under continuous-flow conditions that provide the ultimate test for both catalyst lifetime and immobilization efficiency with regard to activity, selectivity and leaching.12 Immobilization approaches can either use direct metal support through covalent or electrostatic bonding or ligand immobilization by covalent, electrostatic, hydrogen-bonding or aromatic-aromatic interactions (Fig. 2), with the latter being more common. Supports may be metal oxides such as silica, carbon-based materials, polymers, or metal-organic frameworks. Stationary liquid phases such as ionic liquids can also act as supports when containing the solubilized catalyst dispersed on the surface of a highly porous solid.13 Finally, a number of catalyst encapsulation approaches in porous solids seek to trap the molecular catalyst while allowing free passage of the substrate and reaction products.14 When characterizing the hybrid materials obtained after the immobilization step, “classic” characterization techniques of homogeneous molecular organometallic catalysts, such as solution nuclear magnetic resonance (NMR), are not feasible. Nonetheless, solid-state NMR, scanning and transmission electron microscopies (SEM and TEM), besides infrared spectroscopy (IR) and thermogravimetric analysis (TGA) are valuable techniques to assess the immobilized catalyst on the support (qualitatively and quantitatively) and to evaluate the catalyst morphology.

Comprehensive Organometallic Chemistry IV

https://doi.org/10.1016/B978-0-12-820206-7.00071-8

609

610

Separation Strategies in Organometallic Catalysis

Fig. 1 Organometallic catalyst separation strategies.

1.20.2.1

Metal oxide supports

Silica is the most heavily utilized support due to its accessibility, high stability, and porosity.15 Moreover, the Si-OH groups of silica allow robust anchoring of metals or ligands to its surface, generating catalytically active sites. Preparation of Si–O–MXmLn precatalysts is a broad field of research known as surface organometallic catalysis, expanded greatly by the groups of Basset, Marks and Copéret, that has been reviewed extensively over the last decade.16–25 In order to avoid duplication, we will focus our discussion on ligand immobilization.26 Dufaud and co-workers described the covalent immobilization of Wilkinson-type rhodium complexes on mesoporous silica adopting two different strategies aiming to keep the ligand intact.27 In one case, chiral metal complex 1 was tethered to the pore wall of SBA-15 mesoporous silica via an achiral phosphine ligand (1@SBA15, Fig. 3A). In the other approach, achiral metal complex 2 was first embedded into the SBA-3 silica matrix during its synthesis and then chiral auxiliary 3 was grafted into the pores, forming 2/3@PMOS in which the stereogenic and catalytically active centers are held in close proximity within a confined space (Fig. 3B). The performance of the two materials was then assessed for the asymmetric hydrogenation of methyl (Z)-2-Nacetylamidocinnamate in isopropanol. Using 1@SBA-15 (1 mol% Rh), after 72 h under 5 bar H2 at 25  C, 100% conversion was achieved with an enantiomeric excess of 57% for the (R)-enantiomer. Even though phosphine dissociation has been demonstrated in similar catalytic cycles, the performance of 1@SBA-15 almost matched that of the homogeneous catalyst (90% yield with 55% ee using 0.7 mol% Rh). The reusability of the catalyst was evaluated over four cycles without significant loss of activity or enantioselectivity, indicating its robustness. On the other hand, although 2/3@PMOS also afforded the hydrogenation product, both yield (41%) and enantioselectivity (22%) were significantly lower after 120 h at 40  C and 10 bar H2, presumably reflecting the restricted accessibility of the substrate to the embedded Rh catalyst and the heterogeneity of the “chiral pocket.” Thiel et al. also used SBA-15 as a support for the immobilization of catalysts which were applied to the hydrogenation of olefins.28 Their approach consisted of three steps, comprising initial covalent functionalization of the silica with siloxy-substituted imidazolium salts (Fig. 4A) followed by electrostatic grafting of a triarylphosphine ligand (Fig. 4B). Finally, the resulting solid was treated with a divalent Pd precursor, affording the hybrid material (Pd@SBA-15) that was investigated as a catalyst for the hydrogenation of 2-cyclohexenone and other representative alkenes. Pd@SBA-15 was highly active under mild conditions (rt, 1 bar H2, 8–12 h), affording excellent yields of the desired products, and being especially efficient in the presence of double bonds conjugated to an aromatic ring or a heteroatom. Comparison to the homogeneous counterpart was not assessed in this case, although it is known that Pd(PPh3)2Cl2 is a poor catalyst for homogeneous hydrogenation of olefins.29,30 For the hydrogenation of 2-cyclohexenone, the material proved to be highly stable, achieving 99% conversion after up to 10 runs. Characterization of the catalyst using transmission electron microscopy (TEM) strongly suggested the absence of Pd nanoparticles in Pd@SBA-15 even after 10 cycles, suggesting that catalytic activity is due to stable heterogenized molecular catalyst sites. Besides SBA-15, a thorough screening was performed with microporous, mesoporous, and macroporous silica as supports for a 2nd generation Hoveyda-Grubbs catalyst for the heterogeneous ring-opening/ring-closing metathesis of olefins.31 Initial tests were performed in batch reactions in order to assess parameters such as different Ru loadings on the supports and different pore diameters. Through these investigations, the authors found that the most active heterogeneous catalyst for the cyclodimerization of cis-cyclooctene to 1,9-cyclohexadecadiene was the one supported on mesoporous silica MCM-41. Other mesoporous and microporous supports exhibited good conversion but resulted in significant Brønsted acid-catalyzed isomerization of the olefin double bonds. Pore geometry also seemed to affect the selectivity of the catalysts since macroporous supports yielded higher oligomers and polymers which precipitated during the reaction. In the case of MCM-41, it is believed that the Hoveyda-Grubbs complex interacts with the silanol groups on the pores of the support via hydrogen bonding (Fig. 5), providing a balance between activity and

Separation Strategies in Organometallic Catalysis

Fig. 2 Solid support catalyst immobilization strategies.

611

Separation Strategies in Organometallic Catalysis

Fig. 3 Synthesis of (A) 1@SBA15 and (B) 2/3@PMOS.

612

Separation Strategies in Organometallic Catalysis

613

Fig. 4 (A) Immobilized imidazolium chloride on SBA-15 followed by (B) electrostatic grafting of the triarylphosphine ligand.

Fig. 5 Proposed H-bonding of metathesis catalyst in pores of MCM-41.

selectivity. The longevity of this heterogeneous catalyst was evaluated for the same reaction under continuous-flow conditions, reaching up to 8 h with a selectivity of 35% towards the cyclic dimer. Considering the various possibilities of byproduct formation in this cyclodimerization reaction, the authors considered this to be a remarkable result. Silica gel is another versatile choice for covalent immobilization of organometallic catalysts.32–38 In this context, Sawamura et al.32 and Blümel et al.33 adopted the strategy of initially supporting phosphine ligands on the silica and then proceeded to the immobilization of rhodium complexes (Fig. 6A and B). In the former case, the obtained heterogeneous catalysts were studied for the hydrosilylation of ketones with triorganosilanes, showing high tolerance towards sterically hindered ketones and silanes.32 No leaching of Rh into solution was detected and the catalysts could be reused for six cycles with near 100% conversion (total

614 Separation Strategies in Organometallic Catalysis

Fig. 6 (A) Immobilization of Rh complex on ligand-functionalized silica. (B) Formation of Rh-immobilized complex on ligand-functionalized silica. (C) Silica-immobilized Au NHC complexes. (A) Reproduced with permission from Hamasaka, G.; Kawamorita, S.; Ochida, A.; Akiyama, R.; Hara, K.; Fukuoka, A.; Asakura, K.; Chun, W. J.; Ohmiya, H.; Sawamura, M. Organometallics 2008, 27, 6496. (B) Reproduced with permission from Guenther, J.; Reibenspies, J.; Blümel, J. Mol. Catal. 2019, 479, 110630. (C) Reproduced with permission from Sarmiento, J.T.; Suárez-Pantiga, S.; Olmos, A.; Varea, T.; Asensio, G. ACS Catal. 2017, 7, 7147.

Separation Strategies in Organometallic Catalysis

615

TON ¼ 700). The heterogenization of the complex was also likely responsible for the higher reactivity observed for the supported catalysts vs. the homogeneous system: in the hydrosilylation of diisopropyl ketone, for instance, at the same Rh/P ratio (1:1), the heterogeneous catalyst reached 94% conversion in 10 min, whereas the homogeneous system was completely inactive. This behavior can be explained by selective generation of the mono(phosphine)-Rh species on the silica (Fig. 6A), which does not take place under homogeneous conditions. Blümel and co-workers, on the other hand, opted to investigate the catalytic activity of Rh-coordinated phosphinefunctionalized silica (Fig. 6B) in the hydrogenation of 1-dodecene, which reached 100% conversion in 25 h.33 In this case, a change in color (from yellow to brown/black) observed for all the silica-supported materials was the first evidence of a transformation of the catalytically active species over reaction time. Using transmission electron microscopy (TEM), surface-supported Rh nanoparticles were identified on the silica after several runs, indicating decomposition of the Rh complex. Unfortunately, formation of metal nanoparticles is one of the drawbacks when it comes to immobilizing organometallic catalysts, especially under the reducing conditions of hydrogenation reactions.39–41 Nevertheless, catalyst poisoning experiments suggested that the initial catalytically active species was indeed the supported Rh complex, which, after several hours under hydrogenation conditions gradually decomposed, yielding supported Rh nanoparticles. Fortunately, Rh nanoparticles themselves are active catalysts towards hydrogenation40,41 and these immobilized catalysts could be recycled at least 15 times. In turn, Asensio et al. opted to directly immobilize catalysts by grafting siloxy-functionalized gold NHC complexes on silica gel (Fig. 6C).34 The obtained materials were investigated as catalysts in four typical Au-catalyzed nucleophilic addition reactions to alkynes (e.g., hydration of phenylacetylene and hydroamidation of 1-(o-alkynylaryl)urea), presenting high conversions to the desired products. When compared to their homogeneous counterparts, the immobilized complexes exhibited higher activity in some of the reactions tested, such as the cycloisomerization of a 1,6-enyne in wet CH2Cl2, in which the heterogeneous system (1 mol% catalyst) reached complete conversion in 15 min, while the homogeneous analog (5 mol%) required 4 h to convert all the substrate. Repeated batch reactions led to a decrease in their catalytic activity (e.g., complete conversion in the hydroarylation of (prop-2-yn-1-yloxy)benzene took 30 min in the first cycle vs. 10 h in the fourth), although this decrease was not due to leaching. Silica can also be applied in the form of nanoparticles and pellets in order to immobilize organometallic complexes through covalent bonds.42–45 For instance, Li and co-workers described the attachment of a Re(I) carbonyl complex through a disubstituted bipyridine ligand to the surface of silica nanoparticles.42 Since silica nanoparticles are established as suitable solid supports for photocatalysts,46–49 the silica-supported Re complex was able to photocatalytically reduce CO2, reaching comparable activities to the analogous homogeneous system (after 4 h reaction TONheterogeneous ¼ 6.8 and TON homogeneous¼5.8). Using commercial silica pellets, Jacobs and co-workers immobilized a second-generation Hoveyda-Grubbs catalyst.43 The material proved to be efficient towards different types of metathesis reactions, retaining activity over four cycles with full conversion in the ring-opening metathesis polymerization of cis-cyclooctene, corresponding to 400 TON. For the same reaction in a continuous-flow system, the heterogeneous catalyst showed stability for at least 4000 turnovers, one of the highest values reported for immobilized second-generation Hoveyda-Grubbs metathesis catalysts at the time, with very low Ru leaching detected. Aiming to improve catalytic activity and/or to facilitate catalyst separation from the reaction medium, other materials can be incorporated onto the silica support in addition to the organometallic complex. Psaro et al. investigated the immobilization of a cationic Rh complex onto silica-supported palladium nanoparticles (Pd/SiO2), which took place via a dual hydrogen bond/ electrostatic interaction (Fig. 7A).50 This hybrid catalyst was employed in the hydrogenation of benzene to cyclohexane, showing much higher activity as compared to Pd/SiO2, with great stability when submitted to repeated cycles. The authors confirmed enhanced activity only when Rh(I) single sites and Pd nanoparticles are concomitantly anchored to the silica support, suggesting cooperation between the nanoparticles and Rh sites in promoting the hydrogenation of benzene, as illustrated in Fig. 7B. Interestingly, the Pd nanoparticles initially convert benzene to 1,3-cyclohexadiene, while hydrogenation of the latter takes place at the Rh center to irreversibly afford the completely hydrogenated product. Robinson et al. described the reversible electrostatic immobilization of an ammonium-substituted Hoveyda-Grubbs metathesis catalyst on sulfonate-functionalized, silica-coated iron oxide magnetic particles (Fig. 7C).51 This approach allowed for the easy magnetic recovery of the catalyst, simplifying product isolation and catalyst recycling. The obtained heterogeneous catalyst was investigated for the ring-closing metathesis of diethyl diallylmalonate, resulting in >95% conversion to the corresponding cyclopentene after 2 h. However, this hybrid catalyst could be reused for only two cycles before the activity decreased, reaching 56% conversion in the fifth cycle. The loss of activity was shown to result from Ru leaching during the reactions, mainly related to the lability of the benzylidene ligand which generated non-recoverable homogeneous Ru-alkylidenes in the reaction media. Other metal oxides such as alumina, zirconia, and magnesia, besides aluminosilicates such as zeolites, have also been employed as solid supports for molecular organometallic catalysts.6,52–57 Although alumina is a less common support, Scott et al. took advantage of the Lewis acidic sites in g-Al2O3 to attach iridium pincer complexes containing basic and/or polar functional groups.52 Further investigations on the nature of the interaction between the organometallic complex and the support showed that the Ir complexes were covalently anchored to the support via the aryl ring (Fig. 8A).53 The transfer dehydrogenation of cyclooctene to tert-butylethylene in the presence of the heterogeneous catalysts showed high activity towards the desired product with up to 4310 turnovers, with efficient recovery and recycling of catalyst, and very low Ir leaching into solution. Another iridium pincer complex was supported on sulfated zirconia through acetate protonolysis in the work of Delferro et al. (Fig. 8B).56 The hybrid material was tested in the catalytic hydrogenation of propylene and 1,3-butadiene, in both cases showing high activity and selectivity. In the hydrogenation of propylene, no deactivation was observed after 52 h at 100  C whereas for 1,3-butadiene the catalytic activity decreased over time possibly due to coke formation.

616 Separation Strategies in Organometallic Catalysis

Fig. 7 (A) Rh(I) complex [Rh(cod)(dppp)]OTf on silica-supported palladium nanoparticles (Pd/SiO2) (dppp ¼ 1,3-bis(diphenylphosphino)propane) and (B) mechanism proposed for the hydrogenation of benzene by the hybrid catalyst. (C) Quaternary ammonium-substituted Hoveyda-Grubbs catalyst immobilized on sulfonate-functionalized, silica-coated iron oxide magnetic particles. (A and B) Reproduced with permission from Barbaro, P.; Bianchini, C.; Dal Santo, V.; Meli, A.; Moneti, S.; Pirovano, C.; Psaro, R.; Sordelli, L.; Vizza, F. Organometallics 2008, 27, 2810. (C) Reproduced with permission from The Royal Society of Chemistry from Byrnes, M.J.; Hilton, A.M.; Woodward, C.P.; Jackson, W.R.; Robinson, A.J. Green Chem. 2012, 14, 81.

Separation Strategies in Organometallic Catalysis

617

Fig. 8 (A) Proposed structure of the immobilized Ir-complex on g-Al2O3. (B) Grafted Ir-complex on sulfated zirconia.

1.20.2.2

Carbon supports

A plethora of carbon-based materials have been utilized as solid supports for molecular organometallic complexes, including graphene, carbon black, carbon nanotubes, carbon nitrides, and even nanodiamonds.58–67 Typical immobilization strategies include aromatic-aromatic interactions, functionalization of oxidized carbon surfaces (carboxylic acid functional groups), and covalent bonding through diazonium compounds (Fig. 2).68 Messerle and co-workers showed that it was possible to immobilize N,N- and N,P-donor ligands on glassy carbon electrode surfaces through the in situ reductive adsorption of their aryldiazonium salts, resulting in a strong CdC bond between the ligand and the surface. Subsequently, reaction with Rh complex precursors afforded covalently anchored Rh(I) complexes on the glassy carbon.58 The immobilized complexes were investigated as catalysts for the intramolecular hydroamination of 4-pentyl-1-amine to 2-methyl-1-pyrroline. All immobilized Rh(I) complexes bearing N,N-donor ligands were active, yielding only the desired product with 50% conversion, whereas those with the N,P-donor ligand showed low conversion (ca. 15%). TONs reached values of up to 96,000, considerably higher than those of their homogeneous counterparts (TON ¼ 65). Further studies on catalyst stability for three consecutive catalytic runs showed decreasing conversions, indicating that Rh was leaching from the surface during catalysis. X-ray photoelectron spectroscopy (XPS) data for N 1s showed that the ligand itself remained intact, bound to the support, indicating that RhdN bond cleavage may be responsible for the activity decline upon recycling. Several years later, the same research group reported a similar approach for covalent immobilization of a Rh(I) catalyst on high surface area graphene and carbon black.61 Using a carbene-triazole ligand system with a high Rh binding affinity to minimize leaching, this immobilization allowed for efficient reusability of the heterogeneous catalyst. Hydrosilylation of diphenylacetylene with triethylsilane in the presence of the hybrid catalysts reached complete conversion in 15 min at 50  C with carbon black as the support, a similar activity as the homogeneous analog. Recycling experiments showed very good stability over three cycles (24 h each), with complete conversion to the desired product after each run. Even when reused over 10 cycles of 4 h each, complete conversion (>98%) was obtained in all cycles using this Rh(I)-complex immobilized on graphene. Carbon nanotubes have also been employed as solid supports for organometallic complexes through covalent bonding and aromatic-aromatic interactions. Menéndez et al. investigated the functionalization of oxidized and thermally reduced carbon nanotubes with imidazolium salts, which were further used to generate covalently immobilized NHC-carbene Ir-complexes (Fig. 9A).59,62 The hybrid materials were tested as catalysts for the transfer hydrogenation of cyclohexanone to cyclohexanol

Fig. 9 Examples of interactions between organometallic complexes and carbon nanotubes. (A) One of the NHC-Ir complexes covalently anchored on carbon nanotubes. (B) Pyrene-tagged Rh complex immobilized onto carbon nanotubes through p-p stacking. (A) Reproduced with permission from Blanco, M.; Álvarez, P.; Blanco, C.; Jiménez, M.V.; Fernández-Tornos, J.; Pérez-Torrente, J.J.; Oro, L.A.; Menéndez, R. ACS Catal. 2013, 3, 1309. (B) Reproduced with permission from Cunillera, A.; Blanco, C.; Gual, A.; Marinkovic, J.M.; Garcia-Suarez, E.J.; Riisager, A.; Claver, C.; Ruiz, A.; Godard, C. ChemCatChem 2019, 11, 2199.

618

Separation Strategies in Organometallic Catalysis

using 2-propanol as the hydrogen source. Results showed that the Ir complex on partially reduced nanotubes had high catalytic activity, with yields greater than 90% after 100 min. The same behavior was found when the NHC Ir-complexes were attached to oxidized graphene, although the corresponding carbon nanotube-based hybrid catalyst presented the best overall performance.62 In recycling experiments, conversions higher than 90% after five catalytic runs were reported without decreasing Ir content in the reused materials, as assessed by inductively coupled plasma mass spectrometry (ICP-MS) measurements. Moreover, the nanotube-supported catalysts showed air stability, in contrast with the air-sensitive homogeneous Ir–NHC-based catalyst. Godard and co-workers took advantage of p-p stacking interactions to immobilize pyrene-tagged chiral Rh complexes onto multiwalled carbon nanotubes (Fig. 9B).65 The heterogenized catalysts were tested in the asymmetric hydroformylation of norbornene under batch and continuous flow conditions. In batch mode, full conversion of the substrate was achieved at 60  C although with low enantioselectivity (30% ee). At 20  C, 40% conversion was observed with ca. 40% ee. These values are lower when compared to the homogeneous catalyst under the same conditions (62% ee), indicating that the support influenced the enantiofacial discrimination under the reaction conditions. Recycling experiments showed a strong decrease in the conversion after the second cycle, suggesting Rh leaching during hydroformylation. In fact, ICP analysis of the catalyst showed that only 22% of the initial Rh content was still present on the support after the first cycle, with only traces of metal detected after a third run. When the reaction was performed in continuous flow mode, the best conditions at a working pressure of 10 bar and 20  C gave 18% conversion and up to 73% ee. As expected from the results in batch mode, catalyst activity decreased after 100 min in all tests under continuous flow conditions due to metal leaching. This study shows the difficulty typically faced by supported catalysts when strong donor ligands are present in excess. Ruiz-Botella and Peris utilized the same strategy based on p-p stacking interactions to immobilize pyrene-adorned NHC-Rh(I) complexes on reduced graphene oxide.67 Two analogous Rh complexes, one monometallic and the other bimetallic (containing one and two pyrene fragments, respectively), were immobilized on the support and were tested as catalysts for 1,4-addition of arylboronic acids to cyclohexen-2-one and hydrosilylation of terminal alkynes under homogeneous and heterogeneous conditions. In both reactions, the hybrid materials presented similar catalytic activities as their homogeneous counterparts in the first cycle. However, for reusability tests, higher stability and negligible leaching were observed for the bimetallic catalyst holding two pyrene fragments, whereas there was a significant decrease in Rh content for the material containing the monometallic complex with only one pyrene tag. Therefore, these observations point to a dependence of catalyst recyclability on the number of anchoring pyrene sites. For the hydrosilylation of terminal alkynes, the greater immobilization of the bimetallic complex additionally led to a more selective catalyst. In another study, a catechol-phosphine ligand was covalently attached to hydroxyl-terminated nano-diamond support as a strategy to prepare Rh(I)-modified nano-diamonds.69 The phosphine functions on the support were then coordinated to Rh, giving the surface-bound molecular complex. This hybrid material proved to be a selective alkene hydrogenation catalyst in the presence of carbonyl functional groups. After five recycles, no decrease in catalytic activity was observed. Moreover, careful analysis of the solution using inductively coupled plasma atomic emission spectroscopy (ICP-AES) and 31P NMR spectroscopy did not show any evidence of leaching of Rh or ligand P into the reaction medium, respectively. An additional assessment of the solid particles using XPS showed that the P 2p/Rh 3d ratio did not change for the catalyst after five cycles. Carbon nitrides are another class of carbon-based solids that can be used as supports for organometallic complexes, and recent work has especially focused on their photocatalytic applications.60,63 Maeda and co-workers developed functionalized carbon nitrides through the adsorption of phosphonate-functionalized bipyridine ruthenium complexes for the reduction of CO2 to formic acid, with the Ru complex acting as the catalytic unit and the support as the light-harvesting unit (Fig. 10).60 Under visible-light

Fig. 10 Carbon-nitride-supported Ru catalysts for CO2 reduction.

Separation Strategies in Organometallic Catalysis

619

irradiation (l > 400 nm), the selectivity towards formic acid was greater than 80% with the Ru-P complex giving TONs as high as 1000 after 20 h, whereas an analogous homogeneous system under similar conditions gave 68% selectivity to formic acid with 316 turnovers.70 Moreover, the photocatalytic activity of the system deteriorated more quickly under irradiation in homogeneous conditions.

1.20.2.3

Metal-organic frameworks

Recently, metal-organic frameworks (MOFs) bearing organometallic complexes have been used for several original approaches to catalytic reactions.71 Yaghi and co-workers were among the first to describe immobilization of an organometallic complex in MOFs. They prepared a Zn-based MOF in which the organic linkers were functionalized with a Pd(II)-NHC organometallic complex and then assembled into a MOF-type structure.72 In addition to retaining the MOF’s structural order, this immobilization strategy did not significantly diminish its porosity. Porosity retention is one key to good performance of such tailored heterogeneous catalysts since it allows for ready diffusion of substrates to and products from the active sites.73,74 Immobilization of organometallic complexes in MOFs is most often performed through a post-synthetic approach by reacting the metal complex with a previously prepared MOF.75–81 In this case, the catalyst may be immobilized either by covalent bonding to a functionalized linker or encapsulation inside the MOF pores via non-covalent interactions. Lin and co-workers reported the synthesis of Rh- and Ru-functionalized MOFs via post-synthetic metalation of a BINAP-MOF containing the Zr6O4(OH)4(O2CR)12 cluster as a secondary building unit (Fig. 11).75 The Rh-functionalized MOF showed excellent catalytic activity for conjugate addition of arylboronic acids to 2-cyclohexenone and for AlMe3 addition to a,b-unsaturated ketones to yield chiral allylic alcohols. The catalyst gave high enantioselectivity (up to 99% ee) and was three times as active as the homogeneous homolog. In turn, the Ru-MOF was as active as its homogeneous counterpart for the hydrogenation of b-keto esters and substituted alkenes, although with ee’s 2–11% lower. Several tests demonstrated the heterogeneous nature and stability of the MOF-based catalysts, including powder X-ray diffraction (PXRD) of the recovered solids after reactions and ICP-MS of the supernatant that showed minimal leaching of Rh, Ru, and Zr. Another post-synthetic approach was adopted by Rosseinsky et al. in the encapsulation of organometallic cationic Lewis acidic catalyst [CpFe(CO)2(L)]+ (Cp ¼ Z5-C5H5, L ¼ weakly bound solvent) within the pores of an anionic In-based MOF.81 This encapsulation was achieved via a direct one-step cation-exchange, while the MOF retained its porous structure. After thorough characterization of the obtained hybrid MOF material, its activity as a catalyst for the Diels-Alder reaction between isoprene and

Ph Ph P

Ru(MeOH)2Br2

P Ph

Ph

i) Ru(cod)(2-Me-allyl)2; ii) HBr

PPh2 PPh2 [Ru(nbd)2](BF4)

BINAP-MOF

Ph Ph P P Ph

Rh(nbd)+ Ph

Fig. 11 Post-synthetic metalation of BINAP-MOF to obtain Ru- and Rh-functionalized MOFs. Modified with permission from Falkowski, J.M.; Sawano, T.; Zhang, T.; Tsun, G.; Chen, Y.; Lockard, J.V.; Lin, W. J. Am. Chem. Soc. 2014, 136, 5214.

620

Separation Strategies in Organometallic Catalysis

methyl vinyl ketone to 1,4-(para) and 1,5-(meta) isomers of methylacetylcyclohexene was investigated. The heterogeneous catalyst reached 45% yield vs. 67% for the homogeneous counterpart. Nevertheless, the encapsulated catalyst could be separated easily from the reaction medium and reused in two additional cycles. No leaching of iron (from the encapsulated species) or indium (from the MOF) was detected using ICP-OES, demonstrating the stability of the hybrid catalyst. Toste and co-workers used robust IRMOF-10 to achieve the “architectural stabilization” of a homogeneous gold(III) catalyst in order to avoid its decomposition in solution.73 The hybrid material was obtained using a pre-synthetic approach, in which the H2BPDC linker and the gold(III) pre-catalyst Au(III)Cl(H2BPDC)(IPr) were mixed with Zn(NO3)2
4 H2O, affording the Au(III)-MOF (Fig. 12). The catalytically active Au(III) cation in the MOF was accessed by treatment with AgSbF6 and the obtained catalyst was investigated in the cycloisomerization of 1,5-enynes to bicyclohexenes. Substrate conversion reached 44% under heterogeneous conditions with no leaching of Au into solution, demonstrating that the stabilization was effective in preserving the coordinated Au(III) species. Under the same conditions, 68% decomposition of the homogeneous counterpart was observed through detection of Au(I) species in solution. The Au(III)-MOF catalyst showed good recyclability and stability, demonstrating the robustness of the system. In this case, the authors propose that increasing the rigidity of the Au-coordinated biphenyl fragment shuts down the “unimolecular” reductive elimination although mechanistic details were not provided. Another possibility is that an intermolecular mechanism proceeds for the homogeneous case82 that is not possible for the MOF-supported analog. Using a combination of pre- and post-synthetic modifications, Lin et al. incorporated an Ir(III) photoredox catalyst and a Ni(II) cross-coupling catalyst into a stable Zr12-based MOF.83 In this case, the MOF was prepared via solvothermal reaction of ZrCl4 with two mixed linkers, one of them comprising the Ir(III) photosensitizer. After the Zr12-Ir MOF was obtained, post-synthetic metalation with NiCl2
6H2O afforded the Zr12-Ir-Ni MOF, which was tested as a photoredox and organometallic dual catalyst for the cross-coupling of aryl iodides and thiophenols. The reaction between 4-iodobenzotrifluoride and 4-methoxythiophenol in the presence of 0.02 mol% of Zr12-Ir-Ni MOF under irradiation (l ¼ 410 nm) afforded the cross-coupling product in 91% yield with a TON of 4550. In turn, the homogeneous control provided only 5.7% conversion with 2850 TONs. Investigation of the substrate scope of these MOF-catalyzed cross-coupling reactions showed excellent yields for several aryl iodides and aryl thiols, reaching an impressive TON of 38,500 vs. 7.4% yield and 3700 TONs for the homogeneous counterpart. When dealing with reactions in which the optimum conditions involve acidic media, most MOF-based catalysts turn out to be unstable.84 For instance, the best pH for transfer hydrogenation of carbonyl compounds in the presence of homogeneous catalysts was found to be 3.5,85 making the use of MOFs as supports unfeasible when investigating heterogeneous systems. In this case, another class of extended porous frameworks named covalent organic frameworks (COFs), which do not include metal ions in their structure, have also been applied as supports for organometallic complexes. Yoon and co-workers reported Ir and Rh half-sandwich complexes (Fig. 13A) supported on COFs, particularly covalent triazine frameworks (CTF), envisioning that the CTF-metallated heterogeneous catalysts (Fig. 13B) would be active towards the transfer hydrogenation of carbonyl compounds in an aqueous medium.86 The heterogeneous catalysts [CTF(IrCp Cl)Cl] and [CTF(RhCp Cl)Cl] were post-synthetically prepared by reacting the metal complex precursor with a bpy-functionalized CTF (Fig. 13). The resulting catalysts were applied to the transfer hydrogenation of acetophenone, reaching 99% conversion and high selectivity to 1-phenylethanol for [IrCp Cl(CTF)]Cl after 24 h. Other ketones were investigated, and their conversion to the respective alcohol derivatives was also highly selective. While [RhCp Cl(CTF)]Cl proved to be faster (e.g., 99% conversion of acetophenone in 8 h), its stability in recycling experiments was lower with ca. 20% decrease in activity after each run. In contrast, the Ir heterogeneous catalyst could be recycled for at least four runs without any significant loss in activity. The authors believe these

O

O

OH

OH dipp

N

+

+

N

Au Cl

HO

Zn4O(-CO2)6

O

H2BPDC

HO

dipp

O

lPrAu(lll)(H2BPDC)Cl

Fig. 12 Synthesis and structure of Au(III)-MOF (IRMOF-10). Only one gold complex is shown for clarity. Modified with permission from Lee, J.S.; Kapustin, E.A.; Pei, X.; Llopis, S.; Yaghi, O.M.; Toste, F.D. Chem 2020, 6, 145.

Separation Strategies in Organometallic Catalysis

621

Fig. 13 Structural representation of (A) Ir and Rh half-sandwich complexes and (B) CTF-supported complexes. Modified with permission from Sudakar, P.; Gunasekar, G.H.; Baek, I.H.; Yoon, S. Green Chem. 2016, 18, 6457.

results can be related to the reported higher stability of the Ir-hydride intermediate in acidic media, compared to the Rh analog,87,88 although no further discussion of the observed stabilities of the heterogenized catalysts was provided.

1.20.2.4

Polymer supports

Due to their synthetic versatility, polymers are the main choice of organic supports for covalent immobilization of molecular organometallic catalysts.89 As described previously for MOFs, the immobilization process can be performed via pre- or postsynthetic approaches. The former affords main-chain organometallic polymers, in which the transition metal complex is a component of the polymer backbone instead of being supported on the polymer in a post-synthetic strategy.90 Karimi and Akhavan described the synthesis and characterization of N-substituted main-chain organometallic NHC-Pd polymers (NHC-Pd MCOP), aiming to apply them as recyclable catalysts in aqueous media (Fig. 14).91 When tested as catalysts for

Fig. 14 Main-chain NHC-Pd polymers.

622

Separation Strategies in Organometallic Catalysis

Suzuki-Miyaura coupling of 3-bromoacetophenone with phenylboronic acid in water, the NHC-Pd MCOP bearing the more lipophilic n-dodecyl group proved to be more efficient, with 91% yield at low catalyst loading (0.0005 mol%). Similar activity was observed with deactivated and hindered aryl chlorides and bromides, affording good to excellent yields. The catalysts presented high recyclability, although the authors reported the presence of soluble Pd species in the aqueous media (Pd nanoparticles or fragmented NHC-Pd complexes), which are also catalytically active. Ying et al. prepared mesoporous polymer particles (MPPs) using siliceous mesocellular foam microparticles as hard templates, which were further functionalized to support an NHC/pyrrolidine-Pt catalyst.92 The MPPs were prepared from divinylbenzene and 4-vinylbenzyl chloride monomers, followed by immobilization of pyrrolidine via an imidazolium linkage. Subsequently, Karstedt’s catalyst [Pt2(dvtms)3] (dvtms ¼ divinyltetramethylsiloxane) was treated with the functionalized MPPs to generate the NHC-Pt heterogeneous catalyst. In the hydrosilylation of 1-octene, the supported catalyst gave the desired product in 95.5% yield after 22 h at 80  C using a very low Pt loading (0.004 mol%), whereas the homogeneous homolog afforded a similar yield at an even lower Pt loading (0.002 mol%). Nevertheless, the heterogeneous catalyst proved to be as chemoselective (>95%) as the homogeneous counterpart with no platinum leaching to the liquid phase, affording products free of Pt contamination and favoring the recycling of the catalyst. In 2017, Gladysz and co-workers reported their “Catalyst-on-a-Tape” approach to recycling a homogeneous fluorous rhodium catalyst using poly(tetrafluoroethylene) (PTFE) tape.93 This method offers an environmentally friendly alternative to the use of fluorous solvents under liquid/liquid biphasic conditions. The synthesized fluorous Rh-catalyst was investigated in carbonyl hydrosilylation by adding the catalyst (0.15 mol%) to a solution of cyclohexanone and PhMe2SiH in dibutyl ether in the presence of five pieces of PTFE tape. By the end of the reaction, after cooling, the catalyst was supported on the tape, which made possible its easy recovery and recycling (Fig. 15), affording excellent yields over three cycles (98–96% yield, TONs ¼ 666-652). The analogous reaction in the absence of PTFE tape (i.e., homogeneous system) gave the desired product with 98% yield, albeit at higher catalyst loading (1 mol%). The catalyst precoated PTFE tape was also assessed for the hydrosilylation of cyclohexanone, presenting slightly lower yields (94–81%) and TONs (627-543) over three cycles. This work opens interesting perspectives for the use of commercially available fluoropolymers as supports for fluorous catalysts. In a second example, a bis(imidazolium) functionalized polystyrene tag was used in conjunction with (PdCl4)− to generate a robust electrostatically immobilized catalyst system for both batch and flow Heck-Mizoroki coupling reactions (Fig. 16).94 No evidence of imidazolium deprotonation was detected and high TON (5940) and TOF (495 h−1) values were obtained with ca. 50 ppb Pd leaching per cycle. Moreover, the authors also reported a low environmental factor (E-factor), a “green” metric represented by the ratio of the mass of waste obtained per mass of product, using gamma-valerolactone as the solvent. Polymers can additionally assist the retention of homogeneous catalysts in solution in batch or continuous-flow processes. For instance, Vogt et al. reported the use of functionalized polymer particles as phase transfer agents and as catalyst carriers in the hydroformylation of long-chain alkenes.95 The nonpolar cross-linked core of the colloidal polymer particles presented a hydrophilic shell functionalized with cationic groups, which would interact with and attach to any anionic ligand. When added to the multiphasic system for the hydroformylation of 1-octene in the presence of [Rh(acac)(CO)2] as catalyst precursor, the polymer particles were dispersed in the aqueous catalyst phase. This reaction afforded 96% conversion at high substrate:Rh ratios, with TOFs up to 3280 h−1. Recycling of the aqueous catalyst phase was promoted by the polymer particles, avoiding the use of additional organic solvents, and catalytic activity was stable for four consecutive runs with minimum leaching of Rh.

55 °C dibutyl ether

O

–30 °C OSiMe2Ph

reactants +PhMe2SiH

2

fluorous catalyst 3

product

c a taly

P TFE tape

= catalyst

st-c o a te d

cat. coated tape

+2 + PhMe2SiH + fresh solvent 55 °C

decant

3

Fig. 15 Use and recycling of the fluorous Rh-catalyst in hydrosilylation reaction using PTFE tape. Modified with permission from Dinh, L.V.; Jurisch, M.; Fiedler, T.; Gladysz, J.A. ACS Sustain. Chem. Eng. 2017, 5, 10877.

Separation Strategies in Organometallic Catalysis

623

Fig. 16 Polymer-supported Pd/triethylamine catalyst system employed for continuous flow Heck-Mizoroki reaction. Modified with permission from Mahmoudi, H.; Valentini, F.; Ferlin, F.; Bivona, L.A.; Anastasiou, I.; Fusaro, L.; Aprile, C.; Marrocchi, A.; Vaccaro, L. Green Chem. 2019, 21, graphical abstract.

The hydroformylation of 1-octene was investigated by Subramaniam et al. in a continuous-flow reactor in the presence of a Rh-complex containing bulky phosphite ligands bound to a soluble polymer.96 In this case, the reactor was equipped with a polyimide membrane filter that allowed the solvent, products, and unreacted substrate to pass through, retaining the homogeneous polymeric Rh species in solution. Using this continuous-flow system, the hydroformylation of 1-octene reached a constant TOF of approximately 120 h−1 with nearly 50% conversion and 98% selectivity to the aldehyde, which remained constant throughout a 22 h run with insignificant Rh concentrations detected in the effluent. This kind of system could lead to significant savings on industrial processes by reducing precious metal catalyst loss. Porous organic polymers should also be mentioned as a remarkable support option for organometallic catalysts due to their high surface areas and great stability. Jia and co-workers reported the synthesis and characterization of a porous organic polymer via copolymerization of vinyl-functionalized bis(phosphoramidite) and tris(4-vinylphenyl)phosphine ligands, which was further employed as a support for Rh species (Fig. 17).97 The obtained heterogeneous catalyst showed good performance in the hydroformylation of various alkynes, affording the corresponding a,b-unsaturated aldehydes in good to excellent yields with high

Fig. 17 Synthesis of heterogeneous catalyst based on porous organic polymer POL-BINAPa&PPh3 and Rh. Photos reproduced with permission from Liang, Z.; Chen, J.; Chen, X.; Zhang, K.; Lv, J.; Zhao, H.; Zhang, G.; Xie, C.; Zong, L.; Jia, X. Chem. Commun. 2019, 55, electronic supporting information.

624

Separation Strategies in Organometallic Catalysis

Fig. 18 Synthesis of siloxane polymer-encapsulated [Cu/bpy]@Si catalyst (A) employed for atom-transfer radical cyclization of a-halogenated acetamide derivatives (B). Photos reproduced with permission from Motoyama, Y.; Kamo, K.; Yuasa, A.; Nagashima, H. Chem. Commun. 2010, 46, 2256.

selectivity and exhibiting a higher catalytic activity than the comparable homogeneous system. The catalyst could be reused for at least 10 cycles without a substantial decrease in conversion or yield, showing the high stability of this supported Rh catalyst. An additional strategy for immobilizing molecular catalysts consists of encapsulation within porous polymer networks.14,98 In one example, the sequestered [Cu/bpy]@Si catalyst (Fig. 18A) effected the atom-transfer radical cyclization of a-halogenated acetamide derivatives (Fig. 18B) with high turnover numbers (ca. 220 vs. 66 for the homogeneous analog) and without leakage of copper into solution.99 In this reaction, the homogeneous counterpart yielded 66% of the desired product after 1 h, while the [Cu/bpy]@Si heterogeneous catalyst afforded quantitative yields within 4 h. Nevertheless, important achievements of the heterogenization of the catalyst were (i) easy recovery of [Cu/bpy]@Si by simple filtration, facilitating its recycling, and (ii) enhanced air stability after entrapping the catalyst in the polymer, allowing its handling under aerobic conditions for short periods, whereas homogeneous CuCl(bpy) is extremely air sensitive. Additional novel polymer supports include those derived from biomass such as lignin,100 chitin/chitosan,101 cellulose,102 and other polysaccharides.103

1.20.2.5

Supported ionic liquid phase

The use of supported ionic liquid phase (SILP) systems has been increasingly investigated as they allow the use of limited amounts of ionic liquids and present less mass transfer limitations, as opposed to classic biphasic ionic liquid-organic liquid systems (this topic is discussed in Section 1.20.3.2). SILP catalytic systems consist of a thin film of an ionic liquid containing a solubilized metal complex dispersed on the surface of a highly porous solid, usually silica. Wasserscheid and co-workers demonstrated in several publications that SILP catalysts can overcome the issue of mass transport in systems where gas and liquid phases coexist, even allowing the application of fixed-bed reactors in continuous processing when in combination with gaseous reaction mixtures.104–107 For instance, when the hydroformylation of 1-butene was performed in fixed bed reactors using Rh—SulfoXantPhos in [BMIM][n-octylOSO3] (BMIM ¼ 1-n-butyl-3-methylimidazolium) dispersed on silica as the catalyst, the mechanism was the same as for homogeneous Rh-catalyzed hydroformylation.106 TOFs up to 564 h−1 were observed with an exceptionally high average selectivity to n-pentanal under all conditions (ca. 98%). Later, they reported the impact of ligands on the catalyst concentration at the ionic liquid surface in SILP systems typically applied in hydroformylation reactions, showing that the nature of the ligand influences the preferred position of the catalyst complex within or at the surface of a SILP system.108 More specifically, the Rh—tppts complex (tppts ¼ trisulfonated triphenylphosphine) exhibited significant surface activity in the ionic liquid [C2MIM][EtOSO3] (C2MIM ¼ 1-ethyl-3-methylimidazolium), whereas the tppts-free system containing

Separation Strategies in Organometallic Catalysis

625

only the precursor [Rh(acac)(CO)2] showed Rh depletion from the surface. This study allowed a thorough assessment of multiphase homogeneous catalysis using SILP systems, opening the perspective for the manipulation of surface activity of transition-metal complexes in multiphase catalytic systems. In the case of propene hydroformylation, Bell and co-workers investigated the effects of catalyst preparation on the activity and stability of several SILP systems based on Rh complexes containing different phosphine ligands (ca. tppts, XantPhos, and SulfoXantPhos) and ionic liquids (ca. [BMIM][OctSO4], [BMIM][NTf2], and [EMIM][MeOSO3]).109 The assessment of the choice and concentration of ligand and ionic liquid on the catalytic activity showed that the system containing Rh—SulfoXantPhos in [BMIM][OctSO4] was the most active, also exhibiting a higher selectivity towards the n-butanal product (ca. 93% vs. 70–88% for the other systems). Solid state 31P MAS NMR studies indicated that the SulfoXantPhos ligand interacts with the silanol groups in silica (Fig. 19A and B), stabilizing the ligand mainly as species B. Therefore, the presence of sulfonated groups stabilizes the complex in the catalytically active form towards propene hydroformylation, RhH(CO)2(SulfoXantPhos). Interestingly, even the order of addition of the reactants during catalyst preparation can interfere with its activity since the ionic liquid interacts with the silanol groups of the support (Fig. 19C), preventing ligand interaction. Organometallic complexes bearing other metals have also been used in SILP systems. For instance, Wang et al. chose an N-heterocyclic carbene palladium complex (NHC-Pd) and supported it in the mesoporous cage-like material SBA-16 using N-3-(3-trimethoxysilylpropyl)-3-methyl imidazolium chloride as ionic liquid.110 Surface area and pore volume analysis of the obtained material suggested that both NHC-Pd complex and ionic liquid were introduced into the cages of SBA-16, maintaining the mesoporous cage-like structure of the support (Fig. 20). The SILP catalyst was initially tested in the Suzuki-Miyaura reaction using several aryl bromides and phenylboronic acid as substrates, proving to be very active under mild conditions. Complete conversion of bromobenzene within 5 h at 0.01 mol% Pd afforded biphenyl in quantitative yield. For aryl bromides bearing electron-withdrawing or -donating groups, complete conversion was also reached with yields ranging from 87 to 99%. Recyclability tests showed satisfactory yields after nine cycles, although longer reaction times were needed after several cycles (from 5 h in the first

CI

+

N

C N H3 CI −

Pd

CI

N

N

NHC-Pd + lonic liquid

H3 C

Fig. 19 Possible interactions of SulfoXantPhos ligand (A and B) and [BMIM][n-octylOSO3] ionic liquid (C) with silica support.

3C

C H3

N

3

CI

CH

Pd

C N + H3 CI −

H

N

CI N

50 nm

TEM image of SBA-16

cage NHC-Pd/SBA-16-IL

Fig. 20 TEM image of SBA-16 support and schematic description of the preparation of the SILP catalyst containing NHC-Pd complex and ionic liquid. Modified with permission from Yang, H.; Han, X.; Li, G.; Wang, Y. Green Chem. 2009, 11, 1185.

626

Separation Strategies in Organometallic Catalysis

cycle to 49 h in the last). Furthermore, the authors showed the applicability of this catalyst in the Heck reaction using different substrates, obtaining good yields in all cases even after eight cycles of use (ca. 90% yield). A comparison between a biphasic system containing an ionic liquid and 12 SILP systems was carried out by Kukawka and co-workers using Rh complexes for hydrosilylation of 1-octene with 1,1,1,3,5,5,5-heptamethyltrisiloxane (HMTS).111 The SILP systems were prepared with three different Rh-based catalysts and four ionic liquids using calcined silica as support. A considerable difference in activity was observed depending on the ionic liquid, and SILP materials based on ionic liquids with a chemically stable and reactive anion such as [NTf2]− were more active. The most active SILP material carried out 20 cycles with high yields (ca. 71% in the last cycle). Under the conventional liquid-liquid biphasic system, TONs ranged from 31,000 to 91,000, whereas TONs as high as 1,796,000 were reached in the SILP systems. Moreover, SILP conditions allowed the catalyst concentration to be 10 times lower and required only half the reaction time.

1.20.2.6

Concluding remarks

When it comes to catalyst recovery from the reaction media, immobilization on solid supports intuitively appears as a “quick and easy” strategy. However, choice of the most suitable support depends critically on the reaction conditions and potential interactions between the molecular complex and the support. Silica and other metal oxides are well-known, widely accessible supports. In particular, silica presents high stability and porosity, offering the possibility of combination with other materials (e.g., magnetic particles) to facilitate catalyst recovery. For this class of supports, the covalent ligand-support immobilization strategy offers the most robust catalysts with high recyclability. Nevertheless, leaching into the reaction solution is frequently reported to decrease activity upon catalyst reuse, especially in polar solvents or reactions involving strong ligands such as carbonylations. Despite the broad application of silica-supported molecular catalysts in a host of catalytic reactions including olefin hydrogenation and metathesis, hydrosilylation, and photocatalytic reactions, special attention should be paid when dealing with reducing conditions. For instance, hydrogenation reactions can lead to complex decomposition pathways, affording metal nanoparticles that are frequently less active—or even inactive—catalysts for the target transformation. Covalent ligand-support immobilization also appears to be the most suitable strategy for carbon-based supports (e.g., graphene), providing high stability and negligible leaching in hydrosilylation, hydrogenation, and hydroformylation reactions. Immobilization through aromatic-aromatic interactions is another interesting possibility when using graphene and carbon nanotubes as supports. In the case of photocatalytic reactions, carbon nitrides have shown remarkable results and should be considered more widely as potential supports for organometallic catalysts. MOFs are a rising trend in the field of heterogeneous catalysis, offering the possibility to support organometallic catalysts through covalently functionalized linkers or encapsulation in pores. Both strategies can result in negligible catalyst leaching while maintaining the MOF porosity in hydrogenation, Diels-Alder, cycloisomerization, and cross-coupling reactions. For reactions at low pHs, MOFs are less stable and covalent organic frameworks (COFs) may be a better choice. Regarding organic supports, polymers are the main available option, including main-chain organometallic polymers, mesoporous polymer particles, and even PTFE tape. Polymer-supported molecular catalyst have been used successfully for transformations ranging from Suzuki-Miyaura to Heck-Mizoroki coupling and hydroformylation reactions, with negligible complex leaching and good recyclability. Moreover, polymers may be the most suitable option for immobilizing organometallic complexes in continuous-flow processes, opening prospects for their industrial application.

1.20.3

Multiphase strategies for catalyst separation

Multiphase strategies have a long history for efficient separation of molecular catalysts highlighted by the Rührchemie-Rhone-Poulenc aqueous-phase process for propene hydroformylation using a Rh catalyst with sulfonated triphenylphosphine ligands.112 In this strategy, at least two different liquid phases are used to differentially solubilize catalyst and product, enabling the facile separation and recovery of both (Fig. 21).

1.20.3.1

Immobilization in polar and non-polar phases

As typical substrates and products are somewhat soluble in less polar organic solvents, most separation strategies immobilize the organometallic catalyst in a polar phase.113 Particularly, in spite of its potential reactivity, water is an attractive option as the polar phase in biphasic systems since it is non-flammable, non-toxic, non-hazardous, readily available, and easy to separate from most organic solvents.113,114 Furthermore, co-solvents and additives are frequently used to enhance the biphasic interface, thus increasing reaction rates. This has led also to the development of catalysis “on water” after a remarkable acceleration of reaction rates was observed when organic solvents were stirred with water, affording aqueous suspensions.115–117 Further studies on the mechanism proposed that hydrogen-bonding at the water-organic solvent interface plays a key role, with interfacial free OH groups from water molecules establishing strong H-bonds to the transition state rather than the reactants, lowering the activation barrier and enhancing the reaction rate.118,119

Separation Strategies in Organometallic Catalysis

627

Fig. 21 Multiphase strategies for organometallic catalyst immobilization.

In a preliminary investigation, Koridze and co-workers reported the catalytic activity of ferrocene- and ferrocenium-based Pd pincer complexes (Fig. 22) in the Suzuki-Miyaura cross-coupling reaction comparing homogeneous conditions with a multiphase system composed of decane, water, and a liquid phase-transfer catalyst (PTC) tricaprylmethylammonium chloride, A336.120 For the electronically activated substrate 40 -bromoacetophenone, cross-coupling reactions with phenylboronic acid proceeded efficiently under both homogeneous and multiphase conditions using Pd complexes at 0.01 mol%. The multiphase system allowed for the isolation of the diaryl product as crystals from the organic phase in quantitative yields. However, with electronically deactivated substrates (e.g., 4-bromoanisole), the Pd complexes displayed less catalytic activity under multiphase conditions (ca. 60% yield), whereas higher yields of nearly 90% were afforded under homogeneous conditions. One possible explanation would be the formation of a complex between the liquid phase-transfer catalyst A336 and the phenylboronic acid in the multiphase system, which would require a higher concentration of the phase-transfer agent to favor the products,121 although the findings were not conclusive. Nonionic surfactants can also be added to multiphase systems to promote product separation. For instance, Schomäcker et al. tested five different nonionic surfactants in the hydroformylation of 1-dodecene in a multiphase system consisting of water, 1-dodecene, and a surfactant.122 Marlophen NP9, a nonylphenyl ethoxylate, proved to be the most suitable surfactant for the proposed reaction using the hydrophilic catalyst derived from [Rh(acac)(CO)2] and bidentate ligand SulfoXantPhos (Fig. 23).

Fig. 22 Pd pincer catalysts used for multiphase Suzuki-Miyaura cross-coupling.

Fig. 23 [Rh(acac)(CO)2] and SulfoXantPhos complex species (A ¼bisequatorial Rh/SX; B¼equatorial/axial Rh/SX).

628

Separation Strategies in Organometallic Catalysis

In this case, the surfactant enhances olefin solubility at the aqueous/hydrocarbon phase while the SulfoXantPhos ligand generates the water-soluble catalyst that can then be separated from the organic product. A comparison between homogeneous and multiphase systems under similar reaction conditions, with hydrophobic ligand XantPhos used in the former, demonstrated a lower activity for the multiphase system (TOF ¼ 1890 h−1 for homogeneous vs. 295 h−1 for multiphase). Nevertheless, the multiphase system was slightly more selective for the desired linear aldehyde with a 98:2 ratio towards tridecanal, while the homogeneous system afforded a 96:4 ratio. Minimal leaching of rhodium and phosphorous was detected in the product phase using ICP-AES, confirming the applicability of multiphase systems for the hydroformylation of long-chain olefins. Later, Pogrzeba et al. reported the first aqueous, surfactant-free multiphase emulsions as reaction media for the hydroformylation of 1-dodecene using the aforementioned Rh—SulfoXantPhos catalyst (Fig. 23).123 The authors prepared systems containing the olefin, an aqueous solution of the catalyst, and diethylene glycol butyl ether. This short-chain amphiphile could play the role of the surfactant, solubilizing the hydrophobic substrate in aqueous media and facilitating the catalyst recycling. Different reaction parameters were investigated, i.e., amphiphile concentration, syngas pressure, and temperature, and the optimized conditions led to 23.4% conversion after 4 h of reaction at 95  C (TOF ¼ 77.5 h—1). The catalyst was recycled four times, exhibiting high selectivity towards the desired product; however, there was an extensive decrease in the catalyst activity after the first cycle of reuse, which was attributed to amphiphile loss during extraction steps. Nevertheless, the authors demonstrated the potential of aqueous surfactant-free multiphase emulsions as reaction media for homogeneously catalyzed reactions, allowing for an efficient catalyst recycling. When the product of a catalytic reaction is water-soluble the separation strategy must be inverted, and a hydrophobic catalyst is therefore utilized. In 2015, Leitner et al. described the hydrogenation of CO2 to methanol using hydrophobic Ru—Triphos (Triphos ¼ 1,1,1-tris(diphenylphosphinomethyl)ethane) as a homogeneous catalyst, applying a multiphase system for active catalyst recycling.124 The authors employed an aqueous biphasic system in continuous-flow operation, in which the hydrophobic catalyst was retained and recycled in the organic phase (2-methyl tetrahydrofuran; 2-MTHF), whereas product MeOH was removed in the aqueous phase (Fig. 24). The biphasic system presented a slightly lower TON compared to the homogeneous reaction performed in 2-MTHF after 16 h of reaction (247 vs. 258, respectively). However, the catalyst could be recycled three times in the multiphase approach, resulting in a total TON of 769 after four cycles.

1.20.3.2

Immobilization in ionic liquids

Ionic liquids (ILs) offer a tunable liquid phase for the immobilization of homogeneous catalysts, which is able to stabilize ionic transition metal complexes and those that are susceptible to hydrolysis due to their polar, non-aqueous nature. In classic biphasic IL-organic systems, the product is collected in the organic phase while the catalyst is retained in the IL. In 2010, Biffis et al. evaluated the efficiency of bis(carbene) Pd(II) complexes as catalysts under liquid-liquid biphasic conditions using ILs as the catalyst-containing phase for alkyne hydroarylation.125 It was expected that ILs would favor dissociation of anionic ligands from the metal complex, generating the catalytically active species. Moreover, the IL phase would facilitate the dissolution of the strong acid promoter, allowing for facile catalyst separation and recycling. In fact, NTf−2-ILs containing the Pd(II) catalyst substantially increased the reaction rate, achieving yields up to 45% when pentamethylbenzene and ethyl propiolate were utilized, with 54% arene conversion. The reaction performed under the same conditions without IL resulted in a 20% yield with 23% arene conversion. After the first run, the catalyst was efficiently separated from the reaction medium as indicated by analysis of the organic phase using inductively coupled plasma atomic absorption spectroscopy (ICP-AAS), which showed 95% retention of Pd in the IL phase. Recycling of the catalyst-containing phase resulted in a moderate decrease in activity (ca. 10%), although high amounts of the acid co-catalyst were lost to the organic phase due to substantial leaching from the IL phase. A comparison between catalytic activity under homogeneous and biphasic conditions using different ionic liquids was performed by Dupont and co-workers when they investigated the isomerization of estragole to trans-anethole using simple Ru-based complexes.126 Optimized homogeneous conditions achieved 100% conversion with 92% selectivity to trans-anethole,

Fig. 24 Aqueous biphasic system for recycling of the cationic hydrophobic Ru–Triphos catalyst in the organic 2-MTHF phase and removing the product MeOH in the aqueous phase.

Separation Strategies in Organometallic Catalysis

629

using as little as 40 ppm of RuHCl(CO)(PPh3)3. The tests in biphasic systems were initially performed using the IL [BMIM][PF6], which afforded 100% conversion with a 96% selectivity to trans-anethole. However, ICP-AES analyzes detected high Ru content in the organic phase after reactions, indicating that the complex leaches from the IL phase. Addition of an ionophilic phosphine ligand improved the immobilization of the Ru-complex in the IL, although an alternative OH-functionalized IL was needed to increase the activity and selectivity of the catalyst. Unfortunately, even under biphasic conditions, the system could not be reused due to decomposition of the catalytic active species. An advanced strategy relies on the combination of ILs with supercritical CO2 (scCO2), which has been widely investigated as a biphasic process that favors the recovery and recycling of homogeneous organometallic catalysts.127–129 ScCO2 is an interesting alternative to organic solvents since it is highly soluble in the IL phase, being able to extract previously dissolved compounds from the IL phase. The challenge of this process relies on a careful adjustment of the system to enable enough mass transfer between phases, without leaching of IL and catalyst. Walkowiak et al. combined immobilized RuHCl(CO)(PPh3)3 in ILs with scCO2, obtaining a biphasic system for the hydroboration of alkynes using pinacolborane.130 In their approach the catalyst would be retained in the IL phase, while reactants and products would move in and out of the reactor through transport in the mobile supercritical phase. At 100  C, hydroboration of phenylacetylene using 1 mol% of the Ru-complex in ILs afforded yields higher than 90% after 15 min. Low Ru content (ca. 0.07 ppm) was detected in the final products by ICP-MS, showing only slight catalyst leaching. When repetitive batch tests were performed, both conversion and selectivity for hydroboration were stable up to the eighth catalytic cycle. At low temperatures and longer reaction times (e.g., 40  C, 3 h), selectivity to the desired product was even higher. In this example, the use of scCO2 in a biphasic system not only eliminated the need for organic solvents but also favored a significant reduction of the reaction temperature in recyclable hydroboration. Leitner and co-workers demonstrated a biphasic process that allowed for the continuous-flow hydrogenation of CO2 to formic acid, in which scCO2 acted as the mobile phase and an IL was the stationary phase.131 The organometallic catalyst was immobilized in the IL and a stabilizing base was added to produce pure formic acid in a single process unit (Fig. 25). The activity of the IL-adjusted catalytic system containing [Ru(methallyl)2(cod)]/[PBu4][TPPMS] (TPPMS ¼ monosulfonated triphenylphosphine) in [EMIM][NTf2] matched or even surpassed those of similar Ru/PAr3 catalysts in conventional solvents (e.g., benzene), with initial TOFs higher than 314 h−1. Pure formic acid was separated and continuously removed from the system with a scCO2 flow, although catalyst deactivation was observed after 70 h of the experiment as visually indicated by a color change in the system, from bright yellow to orange. Another Ru catalyst system was investigated by Hessel et al. for the methoxycarbonylation of cyclohexene in the presence of an IL and scCO2 in a microflow system.132 After screening different temperatures and pressures, the system comprised of Ru3(CO)12 in [BMIM]Cl under supercritical conditions yielded 77% of the ester product at 180  C and 120 bar, with a residence time of 90 min. Remarkably, the same yield obtained in the supercritical flow system was achieved more than five times faster than that in the subcritical batch system (160  C, 40 bar). Another IL/scCO2 catalytic system was recently reported for the isomerization of allylic alcohols into ketones,133 which interestingly offered the possibility to be combined in a sequential single continuous-flow platform to provide either cyanohydrin or a-amino-nitriles. Initially, the tetravalent Ru catalyst was immobilized in [BMIM][NTf2] for the semi-continuous production of the ketone without the addition of any organic solvent (Fig. 26). Under these conditions, full conversion of allylic alcohol to the corresponding ketone was observed, and another addition of alcohol to the same system gave >95% conversion, reaching TONs up to 1425. Moreover, the colorless scCO2 extracted product did not show any trace of IL or catalyst, whereas the product extracted with Et2O presented a dark-brown color. The feasibility of this isomerization with an immobilized Ru complex in the IL phase, in combination with scCO2, encouraged the authors to further investigate the transformation of the ketone into its cyanohydrin or its

Fig. 25 Process for the direct continuous-flow hydrogenation of CO2 to formic acid based on a two-phase system with scCO2 as extractive mobile phase and an ionic liquid (IL) as stationary phase containing the catalyst and the stabilizing base.

630

Separation Strategies in Organometallic Catalysis

Fig. 26 Schematic representation of the IL/scCO2 setup for the isomerization of 1-octen-3-ol catalyzed by a Ru-complex.

corresponding a-amino-nitrile. The connection of the first system to a second catalytic platform, which depended on the desired product (i.e., cyanohydrin or a-amino-nitrile), afforded excellent yields in both cases (ca. 99%). Interestingly, no purification was required between the catalytic platforms, minimizing waste generation throughout the process.

1.20.3.3

Fluorous biphasic catalysis

Fluorous biphasic catalysis (FBC), a concept first introduced by Horváth and Rábai almost three decades ago, is a notable strategy for heterogenization of molecular organometallic catalysts.134 FBC is based on the use of a liquid perfluorocarbon phase containing the catalyst combined with another phase, which could be a common organic or nonorganic solvent, that has limited or no solubility in the fluorous phase.135,136 In order for the catalyst to be solubilized only or preferentially in the fluorous phase, it should contain enough fluorous moieties (also known as fluorinated ponytails) to favor its solubilization. For instance, Gladysz et al. investigated the phase-transfer activation of a fluorous analog of Grubbs’ second-generation catalyst in ring-opening metathesis polymerization (ROMP).137 Initially, the fluorous analog RuCl2(]CHPh)(SIMes)[P(CH2CH2RF8)3] (SIMes ¼ 1,3-dimesityl-4,5-dihydroimidazol-2-ylidene; RF8 ¼ (CF2)7CF3) was tested as a catalyst in the ROMP of norbornene in CDCl3 solutions, with a slight difference in activity when compared to the nonfluorous Grubbs catalyst RuCl2(]CHPh)(SIMes)(PCy3) (ca. 95 vs. 90% yield after 1 h, respectively). However, when the fluorous solvent perfluoro(methylcyclohexane) (PFMC) was charged with the fluorous analog catalyst, followed by addition of this mixture to a CDCl3 solution of norbornene to create a biphasic system, a remarkable rate acceleration was observed, reaching complete conversion after 10 min. This result can be explained based on the first step of the catalytic cycle, which consists of phosphine ligand dissociation to yield a 14-electron intermediate. In this FBC system, the fluorous phosphine is scavenged by the PFMC phase, allowing the alkene to more effectively compete for the active Ru-catalyst in the organic phase (Fig. 27), contributing to rapid substrate conversion.

Fig. 27 Proposed approach to the phase transfer activation of ROMP of norbornene with a fluorous analog of Grubbs’ second-generation catalyst in a fluorous biphasic system. The black/gray shadings indicate that the active catalyst should have a very high affinity for the reaction phase (CDCl3), while the dissociated ligand has a very high affinity for the fluorous phase (PFMC).

Separation Strategies in Organometallic Catalysis

631

Later, the same research group reported a fluorous/organic phase-transfer activation of ring-closing alkene metathesis using an analog of Grubbs third-generation catalyst with fluorophilic pyridine ligands.138 The reactions were performed with a,o-dienes, including 1,6- and 1,7-dienes, under monophasic and biphasic conditions using the fluorous analog RuCl2(]CHPh)(SIMes) [3,5-NC5H3(CH2CH2RF8)2]2 as catalyst. Reactions of dienes in biphasic systems yielded the ring-closing metathesis products at a faster rate than the monophasic conditions (e.g., 40% conversion in monophasic CD2Cl2 and 80% conversion in the biphasic CD2Cl2/perfluoro(methyldecalin) system, both after 75 min). With this work, the authors extended the generality of phase-transfer activation from phase-tagged phosphine ligands, as previously presented,137 to phase-tagged pyridine ligands. One drawback of conventional fluorous solvents is their high volatility, which restricts their temperature range of application. In order to overcome this issue, the utilization of highly fluorinated ionic liquids combines the efficient extraction property of fluorous solvents with catalyst immobilization in a non-volatile ionic liquid. Rauber et al. reported the synthesis of the highly fluorinated phosphonium ionic liquid [P(RF6)3RF4][NTf2], bearing four perfluoroalkyl chains in the cation combined with the bis(trifluoromethanesulfonyl)imide (NTf−2) anion. The highly fluorinated IL showed limited solubility in organic solvents at room temperature, especially non-oxygenated solvents. Another important characteristic of this IL was its thermomorphic mixing behavior, being reversibly shifted from a biphasic to a monophasic state through a temperature trigger. For instance, with a volume fraction of dimethylformamide (DMF) equal to 0.57, the biphasic system switches to a monophasic state at temperatures above 82  C, reestablishing the two separate phases at lower temperatures. This feature was explored in the Heck cross-coupling of iodobenzene and methyl acrylate or styrene under biphasic conditions (IL/ DMF), utilizing a Pd catalyst containing highly fluorinated ligands, PdCl2[P(m-C6H4-(CH2)2-(CF2)8F)3]2, at 85  C. The yields reached up to 90% with an efficient catalyst recovery achieved for seven runs, shedding light on the potential applications of highly fluorous IL for retaining fluorous catalysts in biphasic systems.

1.20.3.4

Thermomorphic multicomponent solvent systems

Thermomorphic multicomponent solvent (TMS) systems including the one aforementioned have been applied as an additional strategy for the recovery and recycling of homogeneous molecular organometallic catalysts in several reactions,139 not restricted to systems containing fluorous solvents. This approach ensures a homogeneous reaction in a monophasic system at a certain temperature with the recovery of the catalyst through liquid-liquid separation in a biphasic system at lower temperatures (Fig. 28). For instance, the Rh-catalyzed hydroformylation of 1-dodecene was extensively studied in TMS systems comprising DMF and decane.140–142 In these systems, 1-dodecene, the catalyst precursor Rh(acac)(CO)2, and the ligand Biphephos are dissolved in DMF, followed by the addition of the non-polar solvent decane, affording a biphasic system at room temperature. Using this approach, Behr and co-workers performed hydroformylations in a monophasic system at 100  C, achieving yields of 84% towards aldehydes with high n/iso ratios of 98:2.140 Catalyst recycling was performed after cooling down the system, yielding up to 78% of aldehyde in the last of eight cycles. Analysis of catalyst leaching into the product phase using ICP showed low leaching of Rh as well as the P ligand (ca. 7 ppm), indicating a fairly successful application of this TMS system for catalyst recycling. In recent work, Sundmacher et al. recommended the replacement of DMF as catalyst carrier in hydroformylations in TMS systems since it is developmentally toxic.143 A judicious solvent screening indicated diethyl sulfoxide as the most promising solvent, outperforming DMF in terms of catalyst extraction. Nevertheless, the authors acknowledge that further investigations still must be performed to evaluate the hydroformylation reaction performance in the proposed TMS system. TMS systems have also been evaluated by Fischmeister et al. in Ru-catalyzed cross-metathesis of methyl 10-undecenoate with methyl acrylate.144 Reactions were performed using propylene carbonate (PC) as the polar phase and cyclohexane (CH) as the non-polar component. Ethyl acetate (EA) was also added to enable the best phase separation when the reaction mixtures were cooled to room temperature. Optimal reaction performance and phase separation were achieved using a 1:1:2 mixture of PC:EA:CH, reaching 88% of conversion to the desired product, with very low Ru leaching into the product phase as indicated by ICP analysis. When EA was replaced by methyl 10-undecenoate (MU; one of the substrates), cross-metathesis carried out with PC/MU/CH

Fig. 28 Working principle of thermomorphic multicomponent solvent (TMS) systems.

632

Separation Strategies in Organometallic Catalysis

reached 97% conversion in 3 h. Although product extraction and recovery of small polar products were drawbacks of the investigated TMS systems, product recovery was easier with more lipophilic compounds. Water can also be one of the solvents in a TMS system, allowing for a potentially greener process. Vorholt and co-workers described the use of aqueous TMS systems to separate homogeneous transition-metal catalysts from medium-polarity products, employing water as the catalyst phase and 1-butanol as the product phase in the Rh-catalyzed hydroformylation of methyl 10-undecenoate.145 In batch experiments, the water-soluble rhodium catalyst containing the SulfoXantphos ligand yielded 76% of the linear aldehyde after 1 h at 140  C, corresponding to a TOF of 1500 h−1, a remarkable value for the hydroformylation of long-chain olefins in aqueous systems. This TMS system enabled catalyst recycling up to three times with constant yields towards the linear aldehyde. The experiments were later conducted as a continuous process in a mini-plant. Throughout 21 h without interference, a 73% yield was achieved for the linear product, confirming the stability of the catalytic species and the robustness of the TMS system.

1.20.3.5

Concluding remarks

Multiphase strategies for catalyst separation serve as a useful complement to solid-supported approaches. In the simplest versions, ligand modification with polar groups serves to isolate the molecular catalyst from the less polar solvent-containing products with elevated temperatures often serving to ensure productive catalyst-substrate interactions at the interphase boundary. Many sustainable applications employ nontoxic, non-flammable water as the polar solvent although the need for phase-transfer surfactants in some cases necessarily reduces the atom economy. Overall, the low leaching of catalyst to the product phase is a major advantage of these systems. Ionic liquids appear as good solvent alternatives when dealing with water-sensitive organometallic catalysts, with minimal leaching seen for ionic metal complexes. Combining the ionic liquid with supercritical or dense phase CO2 can boost catalyst stability and ease of recovery, as reported above for RuHCl(CO)(PPh3)3. This combination can, however, introduce an additional challenge to the system due to mass transfer issues. The use of a liquid perfluorocarbon phase containing the catalyst is the basis of fluorous biphasic catalysis. In general, this strategy affords the desired products at a faster rate than the monophasic conditions. Nonetheless, the fact that the catalyst must contain fluorous ligands to be soluble in the fluorous phase can be perceived as a limitation to its applicability. The expense and high volatility of the fluorous phase is also a drawback, although the latter can be overcome with the use of highly fluorinated ionic liquids. Finally, thermomorphic multicomponent solvent systems provide an additional and interesting option as a multiphase strategy of catalyst separation. This technique can be applied to temperature-switch systems which are monophasic at one temperature and biphasic system at another, therefore ensuring catalyst recovery through liquid-liquid separation. This strategy is not restricted to fluorinated systems, even offering the possibility of using water as a component. Moreover, very low catalyst leaching has been reported for a number of thermomorphic solvent systems.

1.20.4

Conclusion and outlook

New strategies to promote the separation of molecular organometallic catalysts continue to be developed, often with additional consideration of process sustainability. While simple ligand immobilization approaches can be efficient for reactions involving few potentially coordinating functional groups, for most reactions inevitable ligand dynamics involving substrate- or productcontaining functional groups lead to metal leaching that contributes to reduced catalyst activity or even evolution of soluble catalysts that may exhibit undesired activity and/or selectivity. These factors are especially critical as efforts increase to immobilize first-row metal complex catalysts and employ supported molecular catalysts in continuous flow applications. Multiphasic approaches for organometallic catalyst separation are less susceptible to occasional ligand dissociation and will likely find additional applications as the efficiency of liquid-liquid separations in flow chemistry is improved. With the current drive to develop more sustainable chemical processes and increase computer-driven automation for catalyst synthesis and utilization, we can expect to see additional creative approaches applied to the separation of organometallic catalysts over the ensuing decades.

References 1. 2. 3. 4. 5. 6. 7. 8. 9.

Cornils, B., Herrmann, W. A., Eds.; In Applied Homogeneous Catalysis With Organometallic Compounds, Wiley-VCH Verlag, 1996. Cornils, B., Herrmann, W. A., Eds.; In Applied Homogeneous Catalysis With Organometallic Compounds, Wiley-VCH Verlag, 2002. Cornils, B., Herrmann, W. A., Beller, M., Paciello, R., Eds.; In Applied Homogeneous Catalysis With Organometallic Compounds, Wiley-VCH Verlag, 2017. Cornils, B., Herrmann, W. A., Horváth, I. T., Leitner, W., Mecking, S., Olivier-Bourbigou, H., Vogt, D., Eds.; In Multiphase Homogeneous Catalysis; Wiley-VCH Verlag, 2005. Dupont, J., Kollar, L., Eds.; In Ionic Liquids (ILs) in Organometallic Catalysis; Springer Verlag, 2015. Serna, P.; Gates, B. C. Acc. Chem. Res. 2014, 47 (8), 2612–2620. Benaglia, M., Puglisi, A., Eds.; In Catalyst Immobilization: Methods and Applications; Wiley-VCH Verlag, 2019. Barbaro, P., Liguori, F., Eds.; In Heterogenized Homogeneous Catalysts for Fine Chemicals Production; Springer: Dordrecht, 2010. Hübner, S.; de Vries, J. G.; Farina, V. Adv. Synth. Catal. 2016, 358 (1), 3–25.

Separation Strategies in Organometallic Catalysis

633

10. Pagliaro, M.; Sels, B. F. ChemCatChem 2018, 10 (8), 1663–1665. 11. Mingos, D. M. P., Crabtree, R. H., Eds.; In Comprehensive Organometallic Chemistry III, Elsevier Ltd, 2007. 12. Ishitani, H., Saito, Y., Laroche, B., Rao, X., Kobayashi, S., Luis, S. V., Garcia-Verdugo, E., Eds.; In Flow Chemistry: Integrated Approaches for Practical Applications; Royal Society of Chemistry, 2019; pp 1–49. 13. Romanovsky, B. V.; Tarkhanova, I. G. Russ. Chem. Rev. 2017, 86, 444–458. 14. Otor, H. O.; Steiner, J. B.; García-Sancho, C.; Alba-Rubio, A. C. ACS Catal. 2020, 10 (14), 7630–7656. 15. Thomas, J. M.; Raja, R. J. Organomet. Chem. 2004, 689 (24), 4110–4124. 16. Copéret, C.; Basset, J.-M. Adv. Synth. Catal. 2007, 349 (1–2), 78–92. 17. Basset, J., Psaro, R., Roberto, D., Ugo, R., Eds.; In Modern Surface Organometallic Chemistry; Wiley-VCH Verlag, 2009. 18. Coperet, C. Chem. Rev. 2010, 110 (2), 656–680. 19. Gajan, D.; Copéret, C. New J. Chem. 2011, 35 (11), 2403–2408. 20. Stalzer, M. M.; Delferro, M.; Marks, T. J. Catal. Lett. 2015, 145 (1), 3–14. 21. Copéret, C.; Comas-Vives, A.; Conley, M. P.; Estes, D. P.; Fedorov, A.; Mougel, V.; Nagae, H.; Núñez-Zarur, F.; Zhizhko, P. A. Chem. Rev. 2016, 116 (2), 323–421. 22. Copéret, C.; Estes, D. P.; Larmier, K.; Searles, K. Chem. Rev. 2016, 116 (15), 8463–8505. 23. Coperet, C.; Fedorov, A.; Zhizhko, P. A. Catal. Lett. 2017, 147 (9), 2247–2259. 24. Samantaray, M. K.; Pump, E.; Bendjeriou-Sedjerari, A.; D’Elia, V.; Pelletier, J. D. A.; Guidotti, M.; Psaro, R.; Basset, J.-M. Chem. Soc. Rev. 2018, 47 (22), 8403–8437. 25. Witzke, R. J.; Chapovetsky, A.; Conley, M. P.; Kaphan, D. M.; Delferro, M. ACS Catal. 2020, 10 (20), 11822–11840. 26. Zhong, R.; Lindhorst, A. C.; Groche, F. J.; Kühn, F. E. Chem. Rev. 2017, 117 (3), 1970–2058. 27. Sayah, R.; Framery, E.; Dufaud, V. Green Chem. 2009, 11 (10), 1694–1702. 28. Wang, L.; Dehe, D.; Philippi, T.; Seifert, A.; Ernst, S.; Zhou, Z.; Hartmann, M.; Taylor, R. N. K.; Singh, A. P.; Jia, M.; Thiel, W. R. Catal. Sci. Technol. 2012, 2 (6), 1188–1195. 29. Itatani, H.; Bailar, J. C., Jr. J. Am. Oil Chem. Soc. 1967, 44 (2), 147–151. 30. Natori, I. Polym. J. 2003, 35 (8), 622–627. 31. Bru, M.; Dehn, R.; Teles, J. H.; Deuerlein, S.; Danz, M.; Müller, I. B.; Limbach, M. Chem. Eur. J. 2013, 19 (35), 11661–11671. 32. Hamasaka, G.; Kawamorita, S.; Ochida, A.; Akiyama, R.; Hara, K.; Fukuoka, A.; Asakura, K.; Chun, W. J.; Ohmiya, H.; Sawamura, M. Organometallics 2008, 27 (24), 6495–6506. 33. Guenther, J.; Reibenspies, J.; Blümel, J. Mol. Catal. 2019, 479, 110629. 34. Sarmiento, J. T.; Suárez-Pantiga, S.; Olmos, A.; Varea, T.; Asensio, G. ACS Catal. 2017, 7 (10), 7146–7155. 35. Cook, A. W.; Jones, Z. R.; Wu, G.; Scott, S. L.; Hayton, T. W. J. Am. Chem. Soc. 2018, 140 (1), 394–400. 36. Scott, S. L.; Fu, A.; MacAdams, L. A. Inorg. Chim. Acta 2008, 361 (11), 3315–3321. 37. Nasrallah, H.; Pagnoux, A.; Didier, D.; Magnier, C.; Toupet, L.; Guillot, R.; Crévisy, C.; Mauduit, M.; Schulz, E. Eur. J. Org. Chem. 2014, 2014 (35), 7781–7787. 38. Zelin, J.; Trasarti, A. F.; Apesteguía, C. R. Catal. Commun. 2013, 42, 84–88. 39. Widegren, J. A.; Finke, R. G. J. Mol. Catal. A Chem. 2003, 198 (1), 317–341. 40. Guenther, J.; Reibenspies, J.; Blümel, J. Adv. Synth. Catal. 2011, 353 (2–3), 443–460. 41. Silbernagel, R.; Díaz, A.; Steffensmeier, E.; Clearfield, A.; Blümel, J. J. Mol. Catal. A Chem. 2014, 394, 217–223. 42. Liu, C.; Dubois, K. D.; Louis, M. E.; Vorushilov, A. S.; Li, G. ACS Catal. 2013, 3 (4), 655–662. 43. Van Berlo, B.; Houthoofd, K.; Sels, B. F.; Jacobs, P. A. Adv. Synth. Catal. 2008, 350 (13), 1949–1953. 44. Alesker, M.; Heller, A.; Malik, Z.; Makarovsky, I.; Lellouche, J.-P. J. Mater. Chem. 2011, 21 (29), 10883–10893. 45. Abbott, G.; Brooks, R.; Rosenberg, E.; Terwilliger, M.; Ross, J. B. A.; Ichire, O. O. L. Organometallics 2014, 33 (10), 2467–2478. 46. Gottuso, A.; Köckritz, A.; Saladino, M. L.; Armetta, F.; De Pasquale, C.; Nasillo, G.; Parrino, F. J. Catal. 2020, 391, 202–211. 47. Jadhav, S. A.; Garud, H. B.; Patil, A. H.; Patil, G. D.; Patil, C. R.; Dongale, T. D.; Patil, P. S. Colloid Interface Sci. Commun. 2019, 30, 100181. 48. Chen, X.; Zheng, Z.; Ke, X.; Jaatinen, E.; Xie, T.; Wang, D.; Guo, C.; Zhao, J.; Zhu, H. Green Chem. 2010, 12 (3), 414–419. 49. Badr, Y.; Abd El-Wahed, M. G.; Mahmoud, M. A. J. Hazard. Mater. 2008, 154 (1), 245–253. 50. Barbaro, P.; Bianchini, C.; Dal Santo, V.; Meli, A.; Moneti, S.; Pirovano, C.; Psaro, R.; Sordelli, L.; Vizza, F. Organometallics 2008, 27 (12), 2809–2824. 51. Byrnes, M. J.; Hilton, A. M.; Woodward, C. P.; Jackson, W. R.; Robinson, A. J. Green Chem. 2012, 14 (1), 81–84. 52. Huang, Z.; Brookhart, M.; Goldman, A. S.; Kundu, S.; Ray, A.; Scott, S. L.; Vicente, B. C. Adv. Synth. Catal. 2009, 351 (1–2), 188–206. 53. Vicente, B. C.; Huang, Z.; Brookhart, M.; Goldman, A. S.; Scott, S. L. Dalton Trans. 2011, 40 (16), 4268–4274. 54. Culver, D. B.; Tafazolian, H.; Conley, M. P. Organometallics 2018, 37 (6), 1001–1006. 55. Gu, W.; Stalzer, M. M.; Nicholas, C. P.; Bhattacharyya, A.; Motta, A.; Gallagher, J. R.; Zhang, G.; Miller, J. T.; Kobayashi, T.; Pruski, M.; Delferro, M.; Marks, T. J. J. Am. Chem. Soc. 2015, 137 (21), 6770–6780. 56. Syed, Z. H.; Kaphan, D. M.; Perras, F. A.; Pruski, M.; Ferrandon, M. S.; Wegener, E. C.; Celik, G.; Wen, J.; Liu, C.; Dogan, F.; Goldberg, K. I.; Delferro, M. J. Am. Chem. Soc. 2019, 141 (15), 6325–6337. 57. Campos, J. M.; Lourenço, J. P.; Cramail, H.; Ribeiro, M. R. Prog. Polym. Sci. 2012, 37 (12), 1764–1804. 58. Tregubov, A. A.; Vuong, K. Q.; Luais, E.; Gooding, J. J.; Messerle, B. A. J. Am. Chem. Soc. 2013, 135 (44), 16429–16437. 59. Blanco, M.; Álvarez, P.; Blanco, C.; Jiménez, M. V.; Fernández-Tornos, J.; Pérez-Torrente, J. J.; Oro, L. A.; Menéndez, R. ACS Catal. 2013, 3 (6), 1307–1317. 60. Kuriki, R.; Sekizawa, K.; Ishitani, O.; Maeda, K. Angew. Chem. Int. Ed. 2015, 54 (8), 2406–2409. 61. Wong, C. M.; Walker, D. B.; Soeriyadi, A. H.; Gooding, J. J.; Messerle, B. A. Chem. Sci. 2016, 7 (3), 1996–2004. 62. Blanco, M.; Álvarez, P.; Blanco, C.; Jiménez, M. V.; Fernández-Tornos, J.; Pérez-Torrente, J. J.; Blasco, J.; Subías, G.; Cuartero, V.; Oro, L. A.; Menéndez, R. Carbon 2016, 96, 66–74. 63. Kumar, A.; Kumar, P.; Borkar, R.; Bansiwal, A.; Labhsetwar, N.; Jain, S. L. Carbon 2017, 123, 371–379. 64. Mosconi, D.; Blanco, M.; Gatti, T.; Calvillo, L.; Otyepka, M.; Bakandritsos, A.; Menna, E.; Agnoli, S.; Granozzi, G. Carbon 2019, 143, 318–328. 65. Cunillera, A.; Blanco, C.; Gual, A.; Marinkovic, J. M.; Garcia-Suarez, E. J.; Riisager, A.; Claver, C.; Ruiz, A.; Godard, C. ChemCatChem 2019, 11 (8), 2195–2205. 66. Correa, S. A.; Daniliuc, C. G.; Stark, H. S.; Rojas, R. S. Organometallics 2019, 38, 3327–3337. 67. Ruiz-Botella, S.; Peris, E. ChemCatChem 2018, 10 (8), 1874–1881. 68. Bahr, J. L.; Yang, J.; Kosynkin, D. V.; Bronikowski, M. J.; Smalley, R. E.; Tour, J. M. J. Am. Chem. Soc. 2001, 123 (27), 6536–6542. 69. Queffélec, C.; Schlindwein, S. H.; Gudat, D.; Silvestre, V.; Rodriguez-Zubiri, M.; Fayon, F.; Bujoli, B.; Wang, Q.; Boukherroub, R.; Szunerits, S. ChemCatChem 2017, 9 (3), 432–439. 70. Tamaki, Y.; Morimoto, T.; Koike, K.; Ishitani, O. Proc. Natl. Acad. Sci. 2012, 109 (39), 15673–15678. 71. Dinca, M.; Gabbaï, F. P.; Long, J. R. Organometallics 2019, 38 (18), 3389–3391. 72. Oisaki, K.; Li, Q.; Furukawa, H.; Czaja, A. U.; Yaghi, O. M. J. Am. Chem. Soc. 2010, 132 (27), 9262–9264. 73. Lee, J. S.; Kapustin, E. A.; Pei, X.; Llopis, S.; Yaghi, O. M.; Toste, F. D. Chem 2020, 6 (1), 142–152. 74. Kassie, A. A.; Duan, P.; Gray, M. B.; Schmidt-Rohr, K.; Woodward, P. M.; Wade, C. R. Organometallics 2019, 38 (18), 3419–3428. 75. Falkowski, J. M.; Sawano, T.; Zhang, T.; Tsun, G.; Chen, Y.; Lockard, J. V.; Lin, W. J. Am. Chem. Soc. 2014, 136 (14), 5213–5216. 76. Dau, P. V.; Cohen, S. M. Chem. Commun. 2013, 49 (55), 6128–6130.

634

Separation Strategies in Organometallic Catalysis

77. Syed, Z. H.; Chen, Z.; Idrees, K. B.; Goetjen, T. A.; Wegener, E. C.; Zhang, X.; Chapman, K. W.; Kaphan, D. M.; Delferro, M.; Farha, O. K. Organometallics 2020, 39 (7), 1123–1133. 78. Dau, P. V.; Cohen, S. M. Inorg. Chem. 2015, 54 (7), 3134–3138. 79. Chołuj, A.; Karczykowski, R.; Chmielewski, M. J. Organometallics 2019, 38 (18), 3392–3396. 80. Grigoropoulos, A.; McKay, A. I.; Katsoulidis, A. P.; Davies, R. P.; Haynes, A.; Brammer, L.; Xiao, J.; Weller, A. S.; Rosseinsky, M. J. Angew. Chem. Int. Ed. 2018, 57 (17), 4532–4537. 81. Grigoropoulos, A.; Whitehead, G. F. S.; Perret, N.; Katsoulidis, A. P.; Chadwick, F. M.; Davies, R. P.; Haynes, A.; Brammer, L.; Weller, A. S.; Xiao, J.; Rosseinsky, M. J. Chem. Sci. 2016, 7 (3), 2037–2050. 82. Rocchigiani, L.; Bochmann, M. Chem. Rev. 2021, 121 (14), 8364–8451. 83. Zhu, Y. Y.; Lan, G.; Fan, Y.; Veroneau, S. S.; Song, Y.; Micheroni, D.; Lin, W. Angew. Chem. Int. Ed. 2018, 57 (43), 14090–14094. 84. Mouchaham, G.; Wang, S.; Serre, C. Metal-Organic Frameworks; Wiley-VCH Verlag, 2018. 85. Himeda, Y.; Onozawa-Komatsuzaki, N.; Sugihara, H.; Arakawa, H.; Kasuga, K. J. Mol. Catal. A Chem. 2003, 195 (1), 95–100. 86. Sudakar, P.; Gunasekar, G. H.; Baek, I. H.; Yoon, S. Green Chem. 2016, 18 (24), 6456–6461. 87. Abura, T.; Ogo, S.; Watanabe, Y.; Fukuzumi, S. J. Am. Chem. Soc. 2003, 125 (14), 4149–4154. 88. Pitman, C. L.; Finster, O. N. L.; Miller, A. J. M. Chem. Commun. 2016, 52 (58), 9105–9108. 89. Sommer, W. J.; Weck, M. Coord. Chem. Rev. 2007, 251 (5), 860–873. 90. Williams, K. A.; Boydston, A. J.; Bielawski, C. W. Chem. Soc. Rev. 2007, 36 (5), 729–744. 91. Karimi, B.; Fadavi Akhavan, P. Inorg. Chem. 2011, 50 (13), 6063–6072. 92. Zhang, Y.; Zhao, L.; Patra, P. K.; Ying, J. Y. Adv. Synth. Catal. 2008, 350 (5), 662–666. 93. Dinh, L. V.; Jurisch, M.; Fiedler, T.; Gladysz, J. A. ACS Sustain. Chem. Eng. 2017, 5 (11), 10875–10888. 94. Mahmoudi, H.; Valentini, F.; Ferlin, F.; Bivona, L. A.; Anastasiou, I.; Fusaro, L.; Aprile, C.; Marrocchi, A.; Vaccaro, L. Green Chem. 2019, 21 (2), 355–360. 95. Bibouche, B.; Peral, D.; Stehl, D.; Söderholm, V.; Schomäcker, R.; Von Klitzing, R.; Vogt, D. RSC Adv. 2018, 8 (41), 23332–23338. 96. Fang, J.; Jana, R.; Tunge, J. A.; Subramaniam, B. Appl. Catal. A Gen. 2011, 393 (1–2), 294–301. 97. Liang, Z.; Chen, J.; Chen, X.; Zhang, K.; Lv, J.; Zhao, H.; Zhang, G.; Xie, C.; Zong, L.; Jia, X. Chem. Commun. 2019, 55 (91), 13721–13724. 98. Dergunov, S. A.; Khabiyev, A. T.; Shmakov, S. N.; Kim, M. D.; Ehterami, N.; Weiss, M. C.; Birman, V. B.; Pinkhassik, E. ACS Nano 2016, 10 (12), 11397–11406. 99. Motoyama, Y.; Kamo, K.; Yuasa, A.; Nagashima, H. Chem. Commun. 2010, 46 (13), 2256–2258. 100. Wang, S.; Feng, N.; Zheng, J.; Yoon, K.-B.; Lee, D.; Qu, M.; Zhang, X.; Zhang, H. Polym. Adv. Technol. 2016, 27 (10), 1351–1354. 101. Macquarrie, D. In Heterogenized Homogeneous Catalysts for Fine Chemicals Production; Barbaro, P., Liguori, F., Eds.; Springer: Dordrecht, 2010; pp 1–35. 102. Wolfson, A.; Levy-Ontman, O. Catalysts 2020, 10, 136. 103. Wang, X.; Hu, P.; Xue, F.; Wei, Y. Carbohydr. Polym. 2014, 114, 476–483. 104. Riisager, A.; Fehrmann, R.; Flicker, S.; van Hal, R.; Haumann, M.; Wasserscheid, P. Angew. Chem. Int. Ed. 2005, 44 (5), 815–819. 105. Riisager, A.; Fehrmann, R.; Haumann, M.; Wasserscheid, P. Eur. J. Inorg. Chem. 2006, 2006 (4), 695–706. 106. Haumann, M.; Dentler, K.; Joni, J.; Riisager, A.; Wasserscheid, P. Adv. Synth. Catal. 2007, 349, 425–431. 107. Riisager, A.; Wasserscheid, P.; van Hal, R.; Fehrmann, R. J. Catal. 2003, 219 (2), 452–455. 108. Kolbeck, C.; Paape, N.; Cremer, T.; Schulz, P. S.; Maier, F.; Steinrück, H.-P.; Wasserscheid, P. Chem. Eur. J. 2010, 16, 12083–12087. 109. Shylesh, S.; Hanna, D.; Werner, S.; Bell, A. T. ACS Catal. 2012, 2, 487–493. 110. Yang, H.; Han, X.; Li, G.; Wang, Y. Green Chem. 2009, 11, 1184–1193. 111. Kukawka, R.; Pawlowska-Zygarowicz, A.; Dzialkowska, J.; Pietrowski, M.; Maciejewski, H.; Bica, K.; Smiglak, M. ACS Sustain. Chem. Eng. 2019, 7, 4699–4706. 112. Joó, F.; Papp, É.; Kathó, Á. Top. Catal. 1998, 5 (1), 113–124. 113. Rösler, T.; Faßbach, T. A.; Schrimpf, M.; Vorholt, A. J.; Leitner, W. Ind. Eng. Chem. Res. 2019, 58 (7), 2421–2436. 114. Collis, A. E. C.; Horváth, I. T. Catal. Sci. Technol. 2011, 1 (6), 912–919. 115. Breslow, R.; Maitra, U.; Rideout, D. Tetrahedron Lett. 1983, 24 (18), 1901–1904. 116. Narayan, S.; Muldoon, J.; Finn, M. G.; Fokin, V. V.; Kolb, H. C.; Sharpless, K. B. Angew. Chem. Int. Ed. 2005, 44 (21), 3275–3279. 117. Kitanosono, T.; Masuda, K.; Xu, P.; Kobayashi, S. Chem. Rev. 2018, 118 (2), 679–746. 118. Jung, Y.; Marcus, R. A. J. Am. Chem. Soc. 2007, 129 (17), 5492–5502. 119. Beattie, J. K.; McErlean, C. S. P.; Phippen, C. B. W. Chem. Eur. J. 2010, 16 (30), 8972–8974. 120. Sheloumov, A. M.; Tundo, P.; Dolgushin, F. M.; Koridze, A. A. Eur. J. Inorg. Chem. 2008, 4, 572–576. 121. Paetzold, E.; Oehme, G. J. Mol. Catal. A Chem. 2000, 152 (1), 69–76. 122. Hamerla, T.; Rost, A.; Kasaka, Y.; Schomäcker, R. ChemCatChem 2013, 5, 1854–1862. 123. Pogrzeba, T.; Schmidt, M.; Hohl, L.; Weber, A.; Buchner, G.; Schulz, J.; Schwarze, M.; Kraume, M.; Schoma, R. Ind. Eng. Chem. Res. 2016, 55, 12765–12775. 124. Wesselbaum, S.; Moha, V.; Meuresch, M.; Brosinski, S.; Thenert, K. M.; Kothe, J.; Englert, U.; Hölscher, M.; Klankermayer, J.; Leitner, W. Chem. Sci. 2015, 6, 693–704. 125. Biffis, A.; Gazzola, L.; Tubaro, C.; Basato, M. ChemSusChem 2010, 3 (7), 834–839. 126. Leal, B. C.; Aydos, G. L. P.; Netz, P. A.; Dupont, J. ACS Omega 2017, 2, 1146–1155. 127. Stouten, S. C.; Noël, T.; Wang, Q.; Hessel, V. Chem. Eng. Process. Process Intensif. 2014, 83, 26–32. 128. Gürsel, I. V.; Noël, T.; Wang, Q.; Hessel, V. Green Chem. 2015, 17, 2012–2026. 129. Villa, R.; Alvarez, E.; Porcar, R.; Garcia-Verdugo, E.; Luis, S. V.; Lozano, P. Green Chem. 2019, 21 (24), 6527–6544. 130. Szyling, J.; Franczyk, A.; Stefanowska, K.; Maciejewski, H.; Walkowiak, J. ACS Sustain. Chem. Eng. 2018, 6, 10980–10988. 131. Wesselbaum, S.; Hintermair, U.; Leitner, W. Angew. Chem. Int. Ed. 2012, 51, 8585–8588. 132. Stouten, S. C.; Noël, T.; Wang, Q.; Beller, M.; Hessel, V. Catal. Sci. Technol. 2016, 6, 4712–4717. 133. Peris, E.; Porcar, R.; García-Álvarez, J.; Burguete, M. I.; García-Verdugo, E.; Luis, S. V. ChemSusChem 2019, 12, 1684–1691. 134. Vincent, J.-M.; Contel, M.; Pozzi, G.; Fish, R. H. Coord. Chem. Rev. 2019, 380, 584–599. 135. Horváth, I. T.; Rábai, J. Science 1994, 266 (5182), 72–75. 136. Fish, R. H. Chem. Eur. J. 1999, 5, 1677–1680. 137. Tuba, R.; Costa, R. C.; Bazzi, H. S.; Gladysz, J. A. ACS Catal. 2012, 2, 155–162. 138. Balogh, J.; Hlil, A. R.; Su, H.; Xi, Z.; Bazzi, H. S.; Gladysz, J. A. ChemCatChem 2016, 8, 125–128. 139. Bianga, J.; Künnemann, K. U.; Gaide, T.; Vorholt, A. J.; Seidensticker, T.; Dreimann, J. M.; Vogt, D. Chem. Eur. J. 2019, 25 (50), 11586–11608. 140. Schäfer, E.; Brunsch, Y.; Sadowski, G.; Behr, A. Ind. Eng. Chem. Res. 2012, 51, 10296–10306. 141. Kiedorf, G.; Hoang, D. M.; Müller, A.; Jörke, A.; Markert, J.; Arellano-Garcia, H.; Seidel-Morgenstern, A.; Hamel, C. Chem. Eng. Sci. 2014, 115, 31–48. 142. Dreimann, J. M.; Hoffmann, F.; Skiborowski, M.; Behr, A.; Vorholt, A. J. Ind. Eng. Chem. Res. 2017, 56 (5), 1354–1359. 143. Linke, S.; McBride, K.; Sundmacher, K. ACS Sustain. Chem. Eng. 2020, 8 (29), 10795–10811. 144. Huang, S.; Bilel, H.; Zagrouba, F.; Hamdi, N.; Bruneau, C.; Fischmeister, C. Catal. Commun. 2015, 63, 31–34. 145. Gaide, T.; Dreimann, J. M.; Behr, A.; Vorholt, A. J. Angew. Chem. Int. Ed. 2016, 55, 2924–2928.

1.21

Impurities in Organometallic Catalysis

Nicholas E Leadbeater, Department of Chemistry, University of Connecticut, Storrs, CT, United States © 2022 Elsevier Ltd. All rights reserved.

1.21.1 1.21.2 1.21.3 1.21.3.1 1.21.3.2 1.21.3.3 1.21.4 1.21.4.1 1.21.4.2 1.21.4.3 1.21.4.4 1.21.4.5 1.21.5 1.21.5.1 1.21.5.2 1.21.5.3 1.21.6 References

1.21.1

Introduction Historical perspective Impurities catalyze the reaction Introduction Transition-metal catalyzed cross-coupling reactions Metal-catalyzed cyclization reactions Adventitious contaminants impact the reaction Introduction Catalyst impurities Reagent impurities Solvent impurities The particular case of carbon materials Detecting and avoiding impurities Introduction Analytical tools Approaches to avoiding impurities Concluding remarks

635 635 637 637 637 642 643 643 643 645 646 648 649 649 649 651 653 654

Introduction

“If it looks like a duck, swims like a duck, and quacks like a duck, then it probably is a duck.” The aptly named “duck test” implies that a person can identify an object or phenomenon by observing its habitual characteristics and is sometimes used to counter abstruse arguments that something is not what it appears to be.1,2 In the arena of metal-mediated chemical reactions, this aphorism has come to the fore in recent years. There have been reports of reactions taking place under unprecedented conditions: either mainstream transition-metal catalyzed reactions being performed in the absence of any metal catalyst, or reactions traditionally performed with one particular metal catalyst now working with a very different element. In many of these cases, a thorough assessment of the reaction conditions shows that the reaction takes place via the traditional route, often by the presence of impurities in the mixture.3,4 Well-publicized recent examples include Suzuki coupling reactions taking place by adventitious palladium at the parts-per-billion (ppb) level in an otherwise “palladium free” reaction mixture,5 and silver-mediated click reactions actually being mediated by copper impurities.6 The questions then arise—in the realm of metal catalysis, when is metal-free truly metal-free, and if a reaction traditionally performed using one metal is reported also to work using another, is this due to an impurity of the former in the latter? The aim of this article is to examine these and other key challenges. Following a historical perspective on the topic, a number of case studies will be presented, some general themes highlighted, and some key steps to assessing claims outlined. Overall, experience shows that a duck is indeed a duck, but maybe there are some particular cases when it could be a goose.

1.21.2

Historical perspective

The role of impurities in the outcome of metal-mediated and metal-catalyzed reactions likely goes back to the start of chemistry as a scientific endeavor, but early examples were not reported because the analytical techniques that are at the fingertips of modern chemists were not readily available then. Moving forward to the 20th century, reports started appearing in the 1930s showing that trace amounts of transition-metal contaminants in magnesium can have a profound effect on the preparation of Grignard reagents as well as their subsequent reaction chemistry.7 When the alkyl halide is prone to dimerization (Wurtz coupling) yields of Grignard reagents can vary significantly depending on the impurity profile of the magnesium used (Fig. 1A).8 In the case of bromocyclohexane, a 35% yield of the corresponding Grignard reagent was formed when using European magnesium with a reported 1% of impurities in it. However, when using magnesium from the United States, contaminated with 0.5% aluminum, 0.1% silica, and 0.05% each of iron, manganese, and copper, the yield of the Grignard reagent jumped to 92%. In other cases, transition-metal impurities can have a deleterious effect on the efficacy of the formation of Grignard reagents; for example, they catalyze the reaction of the newly-formed organomagnesium species with unreacted alkyl halide through a radical mechanism.9 Impurities can also change the outcome of a reaction involving a Grignard reagent. When isophorone was treated with methylmagnesium bromide generated from highly-purified magnesium, the 1,2-addition products were formed, but when a trace of copper(I) chloride or elemental copper was present the 1,4-addition products predominated (Fig. 1B).10

Comprehensive Organometallic Chemistry IV

https://doi.org/10.1016/B978-0-12-820206-7.00036-6

635

636

Impurities in Organometallic Catalysis

Fig. 1 Historical examples of metal contaminants.

One example reported in the early 1950s of a metal impurity having a profound effect on the outcome of a reaction indirectly led to a Nobel Prize winning process.11 Ziegler and co-workers were studying a series of polymerization reactions of ethene in the presence of triethylaluminium at elevated pressure.12 On one occasion they observed the formation of 1-butene, the dimerization product of ethene, rather than the usual oligomeric products (Fig. 1C). This anomaly was linked to impurities found in the pressure vessel used for the reactions which had previously been used for hydrogenation of alkenes and traces of nickel remained inside.13 The nickel profoundly changed the outcome of the triethylaluminium-mediated process and in so doing opened avenues for the preparation of butene on an industrial scale but there was a sting to the tail. When the process was first scaled up, it proved unsuccessful. It transpired that the ethene used in the first experiments was contaminated with low levels of acetylene from the manufacturing process. The acetylene proved essential to the success of the process, stabilizing the nickel co-catalyst. Building on these discoveries, Ziegler and his group explored the use of a range of transition-metal additives along with triethylaluminium for alkene polymerization.14 This led to the discovery of a TiCl4 co-catalyzed process that enabled the polymerization of ethene under mild conditions. In 1963, Ziegler, along with Giulio Natta who used the catalysts for stereospecific polymerization reactions, received the Nobel Prize for Chemistry.

Impurities in Organometallic Catalysis

637

Nickel contaminants again came to the fore in the 1960s when the research groups of both Nozaki and Kishi were studying the organochromium-mediated chemoselective addition of allyl, alkenyl, or aryl halides to aldehydes. Independently, the two groups found that the success of the reaction was highly dependent on the source and even the specific batch of chromium(II) chloride starting material.15,16 They subsequently discovered that a trace of nickel(II) chloride in the chromium salt was the culprit; with it the reaction worked well, but without it the reaction failed or competing side reactions predominated (Fig. 1D). The addition, now known as the Nozaki-Hiyama-Kishi reaction, is today performed using CrCl2 but with the deliberate addition of around 0.1 mol% NiCl2. One final historical example, and a case that serves as a bellwether of things to come, relates to the role of lead impurities in zincmediated reactions. Takai, a protégé of Nozaki and the first author of the report on nickel contamination in the chromium-mediated vinyl addition reaction, worked extensively on the development of methodologies focused around the use of organozinc reagents during his independent research career. In the 1980s and 90s, there were instances where reactions that he and his group had developed in Japan involving zinc-carbenoids could not be as effectively performed in the United States or Europe and, likewise, methodologies developed in the United States and Europe were not as successfully reproduced in Japan.17 This prompted Takai to probe in some detail transformations involving zinc as a reducing agent, such as the Simmons-Smith reaction.18 He discovered that traces of lead can have a profound effect on the outcome and that the impurities could be traced back to the process used for isolating and purifying the elemental zinc. When zinc is extracted from its ore and purified by electrolysis of an aqueous solution it is free from lead, but when it is obtained by distillation, it contains around 0.04–0.07 mol% lead. At the time, zinc from Japanese suppliers was generated by the distillation route and that from the United States and Europe came from electrolysis. In the Simmons-Smith reaction, cyclopropanation of cyclooctene with diiodomethane occurred smoothly using European zinc but the reaction did not occur with Japanese zinc (Fig. 1E).18 The lead affects the initial reduction step of diiodomethane by zinc. A similar deactivating effect was observed in a simple reduction of an iodoalkane to an alkylzinc iodide.18 In both cases, the activity can be restored by addition of a small quantity of trimethylsilyl chloride (TMSCl) to the reaction mixture at the start. Surface analysis of zinc generated by electrolysis and distillation showed some significant differences, where the former has a far thinner oxide coating than the latter. One suggestion is that the lead may not actually be involved in the deactivation after all, but more be a bystander in the zinc and that the thicker zinc oxide surface layer may be the root cause of the problem. This would be supported by the fact that TMSCl is reported to abstract oxygen atoms from a metal surface, thereby facilitating the subsequent reduction of the haloalkane reagent.19 However, control studies suggest that the situation might not be so simple. Once a vessel is used for a reaction with leadcontaining zinc, deactivation is also observed in subsequent reactions even when pure zinc is employed, because it is difficult to remove any trace of lead even by washing with aqua regia.17 The issue is further complicated by the observation that lead can actually accelerate closely related classes of reaction. For example, use of distilled zinc, or pure zinc in the presence of a catalytic amount of lead, has a rate-enhancing effect on the methylenation of carbonyl compounds in a Wittig-type olefination process involving diiodomethane, TiCl4 and zinc.20,21 The lead promotes further reduction of the initially-formed zinc carbenoid (ICH2ZnI) to yield a geminal dizinc compound (CH2(ZnI)2), which is a key intermediate in the methylenation process. These cases as a whole show some of the nuances involved when considering the role and impact of impurities on reactions. Overall, in a theme that runs true today, the specific nature of the impurities present can be considerably more significant than the actual quantity. In order to elucidate the exact role of the impurity on the reaction in question, an array of tools can be used. These are discussed later in the article.

1.21.3

Impurities catalyze the reaction

1.21.3.1

Introduction

As history has shown, metal-mediated and metal-catalyzed reactions are not always what they seem at first glance. The role of impurities in a reaction mixture is something that has to be considered. On one hand, while one metal is credited for being a catalyst, the reality is that a second metal, as an impurity in the first, is essentially responsible for the observed outcome. On the other hand, metal-free variants of previously metal-catalyzed reactions have been reported, in which closer examination shows that unexpected metal impurities are the source of the catalytic activity. This has come to the fore in the last decade or two, particularly in the arena of transition-metal catalyzed cross-coupling reactions, as well as some other key transformations. Some of the key examples are highlighted here and summarized in Table 1.

1.21.3.2

Transition-metal catalyzed cross-coupling reactions

An area of catalysis that has been greatly impacted by the issue of impurities is cross-coupling, due not only to its frequent use in academia and industry but also because of the efficiency of these couplings at very low catalyst loadings.39,40 In the late 1990s, there was a flurry of activity focused around use of simple palladium salts as catalysts for the Suzuki coupling reaction (coupling of an organohalide with an organoborane). By employing water as a solvent, it is possible to perform the reaction using low loadings of the palladium salt.41 Nanoparticulate palladium is formed in situ and is responsible for the catalytic activity.42,43 Using microwave irradiation to heat the reaction mixture to above the boiling point of water safely in a sealed tube,44,45 the Suzuki coupling can be complete in as little as 5 min (Fig. 2A).46,47 In 2003, a report emerged of a “catalyst-free” variant of this protocol.22,23 In essence, an aryl bromide, arylboronic acid, and sodium carbonate are placed in water along with tetrabutylammonium bromide as a phase-

638

Table 1

Impurities in Organometallic Catalysis

The impact of impurities on cross-coupling reactions and other key transformations.

Reaction

Original claim

Later discovery

References

Cross-coupling - Suzuki Coupling - Suzuki Coupling - N-, S-, and O-Arylation - Sonogashira Coupling - Sonogashira Coupling - Sonogashira Coupling - Grignard reagents with aryl halides Azide-alkyne cycloaddition Heterocycloisomerization

Metal-free Iron-catalyzed Iron-catalyzed Copper-catalyzed Iron-catalyzed Gold-catalyzed Manganese-catalyzed Silver-catalyzed Metal-free

Catalyzed by trace palladium Catalyzed by trace palladium Catalyzed by trace copper Catalyzed by trace palladium Catalyzed by copper or trace palladium Possibly catalyzed by trace palladium Catalyzed by trace palladium or copper Catalyzed by trace copper Catalyzed by trace copper

5,22,23 24–26 27 28 29 30–34 35,36 37 38

(A)

(B)

(C)

Fig. 2 Palladium contaminants in “Metal-Free” Suzuki coupling reactions.

transfer agent. The reaction mixture is heated at 150  C for 5 min in a sealed tube after which time the biaryl is obtained in good to excellent yield (Fig. 2B). Care was taken to ensure that obvious potential sources of palladium were avoided; new reagent bottles, new glassware and spatulas, and triply distilled water were employed. In addition, assay of the reaction mixture did not show evidence of palladium down to the level of detection of 1 ppm. This added credence to the proposition that the reaction was “metalfree.” However, over time, issues started emerging. The group, based in the United Kingdom, first found that the work could not be easily repeated in Sweden where they were attempting to perform a series of scale-up studies. This coincided with the group moving permanently to the United States, where again the work could not be reproduced. This led to the team critically assessing the methodology. They discovered that palladium contaminants at the ppb level were responsible for the coupling reaction. Indeed, the addition of as little as 50 ppb of palladium to an otherwise “metal-free” reaction mixture was sufficient for the coupling to be complete within a five-minute timeframe (Fig. 2C).5,48 It was later discovered that sodium carbonate from the United Kingdom contained ppb levels of palladium, whereas that from Mainland Europe and the United States did not. While this was the first report of a cross-coupling reaction catalyzed by an impurity, it was not the last. Palladium salts and complexes are costly and also toxic. In efforts to overcome these issues, the design and use of nonprecious metals as catalysts for cross-coupling reactions has come to the fore.49 Of these, iron complexes started receiving particular attention in the late 2000s.50 With iron being abundant, cheap, and having a better safety profile, it seemed that this line of research could be very impactful. In 2008, two reports emerged of iron-catalyzed Suzuki coupling reactions. One involved use of a ferrocene-based complex51 and the other entailed pyridine-ligated iron complexes.24 In the case of the latter, the reaction was performed at 80  C for

Impurities in Organometallic Catalysis

639

Fig. 3 Palladium contaminants in an “Iron-Catalyzed” Suzuki coupling reaction.

24 h in an aqueous solvent mixture using 1 mol% of Fe(py)4Cl2. The use of tetrabutylammonium bromide enhanced the reaction. This observation, along with the fact that the coupling is performed in water using a simple iron complex, suggests an analogy to the problematic chemistry above. The conditions are very similar to those previously reported for the catalyst-free coupling, and it could feasibly be that in this case the iron complex was just a bystander and that the catalytic activity was as a result of palladium contaminants (Fig. 3). Another research group probed the reaction in more detail and found it not to be reproducible, but addition of as little as 0.1 ppm palladium to the mixture then mirrored the catalytic activity reported.25 This clearly indicates that trace palladium in either the reagents, glassware, or solvent was responsible. The original article was subsequently retracted.26 Less than a year later another report emerged,52 followed by a retraction.53 This time, iron(III) chloride was touted as a catalyst, when used in conjunction with ethanol as the solvent and potassium fluoride as a base. The reaction mixture was heated at 100  C for 16 h in a sealed tube. Shortly after being published, issues emerged; the methodology could not be reproduced by other groups.27 On this occasion, the potassium fluoride was to blame (Fig. 3). Further study showed that the batch of base used in the development work was most likely contaminated with trace metals and this was the origin of the catalytic activity.53 The reaction did not work with other sources of potassium fluoride. In addition to these clear-cut examples of impurities being the root of the catalytic activity, other questionable reports have emerged.54,55 In these, reproducibility remains an issue.56,57 In one case this precipitated retraction by the authors58 and in others evidence points to either trace impurities in the reaction mixture, or mischaracterization of the iron species responsible for catalytic activity. Overall, iron-catalyzed Suzuki coupling may well be possible, but probably remain limited to either highly reactive heterocyclic substrates or aryl halide substrates that can participate in directed activation.59 Reports of iron complexes as purported catalysts for cross-coupling goes beyond just forging carbon-carbon bonds, extending to CdO and CdN bond forming reactions (Fig. 4).60 Simple iron salts like iron(III) chloride, in conjunction with a diamine or diketone ligand, have been used to prepare arylated phenols,61 thiols,62 sulfoximines,63 and amines,64 as well as nitrogen heterocycles65,66 and amides.67 However, the catalytic activity was dependent on the purity of the FeCl3 used; higher purity led to a lower yield of the coupled product. In addition, even the source of the iron(III) chloride impacted the reaction. This led to a reassessment of the methodologies.27 In this case, instead of palladium, copper was found to be the culprit. When using iron(III) chloride of greater than 99.99% purity, addition of as little as 10 ppm of copper(I) oxide led to a significant increase in yield. In two cases studied, essentially identical results could be obtained if the reaction was performed using 10 ppm Cu2O but no FeCl3.27,68 The role of impurities in cross-coupling reactions has also come to the fore in the Sonogashira reaction. This coupling of an aryl halide with an alkyne is traditionally performed with either a palladium catalyst or a mixture of palladium and copper species.69 In the case of the latter, the role of the copper is to provide a source of the requisite acetylide species for the transmetallation step. That said, with the increasing interest in employing cheaper, less toxic metals as catalysts for reactions, a range of “copper-only” Sonogashira coupling protocols have emerged.70 Discrete copper complexes, combinations of a copper salt and a ligand, as well as copper nanoparticles have all been featured. However, traces of palladium, down to the ppb level, dramatically enhance the rate of reaction (Fig. 5A).28 Such small amounts of palladium could feasibly originate from the copper source, the reagents, the base, or the solvent in the reaction mixture. The reaction was effective even at palladium concentrations of 10 ppb. It could well be that the catalysis is palladium-driven, and that the copper simply serves as a source of acetylide.

Fig. 4 Copper contaminants in an “Iron-Catalyzed” CdO and CdN coupling reactions.

640

Impurities in Organometallic Catalysis

(A)

(B)

Fig. 5 Palladium contaminants in Sonogashira coupling reactions.

In 2003, two reports of “transition-metal free” Sonogashira couplings appeared in quick succession (Fig. 5B).71,72 In many regards they mirrored the protocol for the “metal-free” Suzuki coupling in that they involved the use of water as a solvent, a phase transfer agent, and elevated temperatures.22,23 Sodium hydroxide is added as a base in one case,71 and sodium carbonate in the other.72 The scope of the protocols is limited to aryl bromides as the organohalide component and phenylacetylene as the alkyne coupling partner. Given that parts per billion levels of palladium are enough to catalyze the Sonogashira reaction,28 it is entirely possible that these “metal-free” couplings are in fact palladium-catalyzed, the metal coming from one or more components of the reaction mixture or the apparatus used.30 A number of other catalyst-free variants have appeared in the literature since these first reports,73 and again the claims in these methodologies need to be considered cautiously. Issues with iron(III) chloride have also come to the fore in the arena of Sonogashira coupling reactions. In the presence of a diamine ligand and a suitable base, aryl iodides could be coupled with a range of alkynes.74 Iron(III) acetylacetonate has also been reported to catalyze the reaction.75 Since adventitious copper was responsible for the catalytic activity in FeCl3-mediated C-N, C-O, and C-S couplings, it was only logical that the Sonogashira reaction should be reassessed, uncovering a problem.29 In the presence of a large excess of diamine ligand, the reaction was effectively catalyzed by very low concentrations of copper. But then the nagging question is whether the reaction really is copper-catalyzed. Bearing in mind the observation that ppb levels of palladium are enough for high catalytic activity,28 could it be that the reaction is actually palladium-catalyzed? While the copper catalyst was analyzed for other metals and only 4 ppb of palladium was found, the reagents, solvent, reaction vessel, and stir bar appear not to have received the same scrutiny. If any of these brought palladium into the reaction mixture, there may be enough present to account for the catalytic coupling activity. Despite all these red flags, “iron-catalyzed” variants of the Sonogashira reaction are still appearing in the literature, without the requisite tests for the presence of palladium.76 Other metals have been reported as catalysts for the Sonogashira reaction, including cobalt,77 ruthenium,78 nickel,79 silver,80 gold,31,81,82 and indium.83 Not unexpectedly, controversy and even arguments have ensued, none more so than in the case of gold “catalysts.” In 2008, gold supported on nanocrystalline cerium oxide was reported to catalyze the Sonogashira reaction between iodobenzene and phenylacetylene (Fig. 6A).31 The catalytically active species was suggested to be gold(I), which has the same electron configuration as elemental palladium and copper(I). Shortly afterwards, other reports emerged suggesting that gold(I) can catalyze the reaction.82,84 However, doubt was cast over these results.32 To start with, the coordinative preferences of Au(I), Pd(0), and Cu(I) are not very similar despite their common d10 configuration. Secondly, oxidative addition, the first step of the reaction, does not occur in a stoichiometric reaction (Fig. 6B), which suggests that gold(I) should not be able to catalyze the Sonogashira reaction. Finally, there were issues with reproducibility. Given that commercial gold is often contaminated by traces of palladium, the conclusion drawn was one that is the ongoing theme in this article: the reaction is catalyzed by adventitious palladium, be that from the gold or the reagents, solvent, or glassware. The story does not end here though. In order to determine whether gold supported on cerium oxide is able to catalyze the Sonogashira reaction or if palladium impurities are indeed the key species in this process, the authors of the initial report performed a series of kinetic studies of the coupling between iodobenzene and phenylacetylene.33 Keeping the amount of Au/CeO2 constant, they performed the reaction at different concentrations of palladium. The premise was that if palladium is solely responsible for the catalytic activity, the initial rate constant for the reaction would approach zero as the concentration of palladium was decreased. This is not what was experimentally observed, suggesting that there is some intrinsic activity from the gold nanoparticles. This assertion is questionable though. Even when using the reported 99.999% pure gold, the remaining 0.001% still counts for at least 10 ppm. If some or all of this is palladium, the concentration is still orders of magnitude greater than that already discussed in this article for catalysis of coupling

Impurities in Organometallic Catalysis

641

(A)

(B)

Fig. 6 Palladium contaminants in “Gold-Catalyzed” Sonogashira coupling reactions.

reactions. The same logic could be applied to another series of reports by surface chemists suggesting that gold can catalyze the Sonogashira coupling. Phenylacetylene and iodobenzene react on a smooth Au(111) surface under vacuum conditions to yield biphenyl and 1,4-diphenylbutatdiyne together with the Sonogashira coupling product,85 and nanoparticulate gold supported on CeO2 does show selectivity for the cross-coupling product.86 While the debate continues in the literature,87 it would be fair to say that the jury is still out regarding whether gold does or does not catalyze the Sonogashira reaction, or indeed other coupling processes.34,88 More study is required before a definitive answer can be reached. The Heck reaction serves as yet another case in point when it comes to the potential role of impurities as a source of catalytic activity in cross-coupling processes. The coupling of an organohalide with an alkene is traditionally performed with a palladium catalyst.89 Over time, conditions have been developed whereby the Heck reaction can be affected using ultra-low (sometimes termed “homeopathic”) levels of palladium as the catalyst.90–92 Again, water turns out to be an excellent solvent for performing the reaction at low catalyst loadings.93 Another offshoot of the protocol for a “metal-free” Suzuki reaction was the report of a catalystfree Heck coupling in supercritical water.94 Using potassium acetate as a base, iodobenzene and styrene were reportedly coupled to give stilbene after 10 min at 375  C and 250 bar (Fig. 7A). Although a feasible mechanism has been put forward, it is not unreasonable to believe that residual palladium in either the base used for the reactions or in the reactor vessels themselves may be the catalyst. While product mixtures were analyzed for Ni, Fe, Cr and Mo, no mention is made of palladium. A reassessment of the methodology shows that as little as 500 ppb of palladium is enough for the reaction to proceed efficiently (Fig. 7B).95 Despite there being doubt over the verisimilitude of palladium-free Heck couplings, reports continue to emerge using other metals. Examples of copper,96 iron,97 and cobalt98 complexes as “catalysts” have appeared in the literature. Given that the necessary steps have not been taken to ensure that traces of palladium are not present, these reports need to be approached with skepticism. A final example of a cross-coupling reaction actually shows new chemical space and impurities within the same manifold. For some time, manganese has been touted as a cheaper, less-toxic alternative to palladium as a catalyst for the reaction of organohalides with suitable coupling partners.99 Of these, the coupling of aryl halides with Grignard reagents using manganese(II) chloride (A)

(B)

Fig. 7 Palladium contaminants in “Metal-Free” Heck coupling reactions.

642

Impurities in Organometallic Catalysis

(A)

(B)

Fig. 8 Metal contaminants in “Manganese-Catalyzed” coupling reactions.

has received significant interest.35,100,101 It is possible to perform the reaction quickly, easily, and effectively. In a couple of these cases, all care and attention was taken to ensure that contamination from other metals was not playing a role.35,36 Mechanistic studies suggest a radical mechanism, which is well-precedented in manganese chemistry. The development of a manganesecatalyzed approach for the coupling reaction does therefore have credence. However, older reports of manganese-catalyzed reactions have started to receive more scrutiny. Two such examples are the reaction of aryl halides with organostananes36 and with amines,102 which bring about CdC and CdN bond-forming reactions respectively. In these cases, a radical pathway is less likely. A critical reassessment of these methodologies shows that the manganese salt that is added to the reaction is not playing a role and that the catalytic activity comes from another metal.36 For the CdC coupling, traces of palladium are responsible (Fig. 8A) and for the CdN coupling, adventitious copper is the catalytically active species (Fig. 8B).

1.21.3.3

Metal-catalyzed cyclization reactions

When discussing the role of impurities in metal-catalyzed reactions, two particular examples of cyclization reactions deserve a mention. The first involves the now-ubiquitous “click” reaction, the cycloaddition of an azide with an alkyne to generate a 1,2, 3-triazole.103 The reaction can proceed uncatalyzed (known as the Huisgen 1,3-dipolar cyclization) but requires elevated temperatures and often produces mixtures of two regioisomers of the triazole product when using non-symmetric alkynes. When using copper(I) salts as catalysts, the reaction can be performed in aqueous solution at room temperature, and yields only the 1,4-disubstituted 1,2,3-triazole regioisomer.104 With an eye to improving the methodology, and also accessing the complementary, 1,5-disubstituted regioisomer, chemists have examined metals other than copper as catalysts.105 Of these, silver complexes have received particular attention.106 Simple silver(I) salts prove not to be particularly effective, but when combined with ligands, or if pre-formed silver(I) complexes are used, the catalytic activity increases considerably.107 One such example is a dimeric silver complex bearing the tripodal triazole ligand, tris(benzyltriazolylmethyl)amine (TBTA), [Ag2(TBTA)2](BF4)2(Fig. 9).6 However, when using this catalyst, yields of the triazole products vary considerably. Even when using crystals of the silver complexes that are of sufficient quality for structure determination, and confirming consistent molecular structure from different preparations was confirmed by X-ray crystallography, reproducibility in the cycloaddition persists. The question arises as to whether the catalytic activity is truly due to the silver complex. Given the chemical similarity between copper and silver, the fact that the TBTA ligand is formed by means of a copper-catalyzed click reaction and is itself a good ligand for copper, the possibility of copper contamination has to be considered. Indeed, crystalline samples of [Ag2(TBTA)2](BF4)2 were found to contain around 560 ppm of copper when reagent grade silver tetrafluoroborate (>98% purity) is used to prepare the complex. Even when using an alternative purification process for the ligand, approximately 200 ppm of copper is still present; likely enough to catalyze the click reaction. After changing to high-purity AgBF4 ( 99.99% purity), the copper contamination drops to a level of around 8 ppm, at which point no catalytic activity is observed. This example drives home the insidious nature of impurities. Even when using ultrapure material, traces of other metals can pose a challenge. It also means that examples of click reaction purportedly catalyzed by silver, which continue to appear in the literature,6,37 need to be considered with extra care and attention.

Impurities in Organometallic Catalysis

643

Fig. 9 Copper contaminants in a “Silver-Catalyzed” click reaction.

Fig. 10 Impact of an impure starting material on a cyclization reaction.

The second example of a cyclization process that turns out to have a twist to the tale involves the heterocycloisomerization of alkyne-functionalized substrates. These reactions are traditionally performed using p-Lewis acid catalysts based on platinum, copper, or silver.108 However, in one recent instance the heterocycloisomerization reaction was reported to proceed simply by heating the substrate in water or methanol.109 In addition, a higher yield of the desired product was obtained in the absence of a metal “catalyst.” Another example of a “no added metal catalyst” reaction was then found, namely the heterocycloisomerization of a series of alkynylcyclopropylhydrazones to yield cycloheptane-fused aminopyrroles.110 This observation warranted further investigation, the results of which showed that the process was not metal-free. Traces of copper contaminants in the reaction mixture were responsible for the cyclization (Fig. 10). The origin of the copper could be traced back to a Sonogashira reaction used to install the alkyne functional group in the cyclization substrate. Of note is that this coupling was performed three synthetic steps before the final substrate was obtained, each of these steps having a standard purification (extraction or column chromatography). This shows the pernicious nature of metal contaminants. Another noteworthy facet of this work was the combination of computational and experimental studies to determine the true mechanism of the cyclization reaction. Such an approach can yield insight which could not have been obtained with such a high degree of certainty by utilizing only one of the two approaches.38 In this case, computational work elucidated a feasible mechanism, and experimental studies using in-situ reaction monitoring and elemental analysis allowed for positive identification of key reaction intermediates along the mechanistic pathway.

1.21.4

Adventitious contaminants impact the reaction

1.21.4.1

Introduction

In addition to cases where it is solely responsible for the observed catalytic activity, there are examples where an adventitious contaminant can enhance the efficacy of an already effective process. The impurity plays a role in the reaction, either prolonging the lifetime of the catalyst, increasing its turnover number or turnover frequency, or facilitating the reaction through a secondary pathway. Impurities can also have a negative impact on the outcome, either deactivating the catalyst or opening avenues for competitive, off-target pathways. To illustrate this, a number of key recent cases in point are discussed, and are summarized in Table 2.

1.21.4.2

Catalyst impurities

Of all the palladium complexes that are used as catalysts or catalyst precursors, palladium(II) acetate is one of the most popular (Fig. 11). It is employed both commercially and in academic settings in a wide range of bond-forging reactions itself, as well as serving as the starting material to more elaborate catalysts.118 Palladium(II) acetate is a trimeric complex, Pd3(OAc)6 (1), but

644

Table 2

Impurities in Organometallic Catalysis

The impact of adventitious contaminants on key reactions.

Reaction

Adventitious component

Observation

References

Catalysis using palladium acetate Ring-opening metathesis polymerization

Two palladium impurities Alkyne-functionalized norbornene

Impurities impact reactions differently Catalyst deactivated by contaminant from starting material synthesis Linear rather than cyclic polybutadienes formed Contaminants improve catalytic performance

111 112

Contaminant from solvent enhances catalytic activity Process enhanced by contaminant from catalyst synthesis step Catalytic activity can be enhanced, reduced, or completely inhibited

115 116 117

Ring-closing metathesis Ruthenium-catalyzed arylation Photocatalytic hydrogen evolution Metal-mediated Michael addition

4-Vinylcyclohexene Impurities in diethyl diallylmalonate g-Butyrolactone Silver salts Halide impurities in ionic liquids

Fig. 11 Impact of impurities on the catalytic activity of palladium(II) acetate.

113 114

Impurities in Organometallic Catalysis

645

commercially available sources are often contaminated by two specific compounds.119 The first, Pd3(OAc)5(NO2) (2), has one acetate ligand replaced by a nitrite moiety, a stronger bidentate ligand.120,121 It is an artifact of the preparative route to palladium acetate which involves the treatment of metallic palladium with a mixture of nitric and acetic acids. The other component is a polymeric species, [Pd(OAc)2]n (3) the quantity of which is dependent on the amount of water present in the preparative route.120 While polymers are generally present only in trace amounts, commercial lots of palladium acetate can contain up to 20% of the nitrite-ligated impurity.120 In addition, the parent compound 1 can slowly convert to hydrolysis product 4 if left open to the air; up to 34% of 1 can turn to 4 in a week.122 When it comes to catalytic activity, in some cases the impurities do not have a sizable impact, but in others they do. The outcome is dependent on factors such as the nature of the coupling, the particular substrate involved, and the temperature at which the reaction is performed.123,124 At elevated temperature and in the presence of a bidentate phosphine ligand, prototypical Suzuki couplings (CdC bond formation) and Buchwald-Hartwig aminations (CdN bond formation) seem relatively unperturbed by the nature of the catalyst, be it 1, 2, or 3. In the case of polymeric palladium complex 3 in organic solvents, the ligand may play a key role in generating a soluble palladium species from what is essentially an insoluble precursor. If the same Suzuki reaction is performed at room temperature, there is a marked diminution in the activity of 3. This could well be a matter of solubility as well. This same theme is true when it comes to a-carbonyl arylation. In the absence of an added ligand, 1 and 2 work well but 3 performs poorly, this time regardless of temperature. Adding a phosphine to the reaction mixture levels the playing field; 3 is just as active as its trimeric congeners.124 In Heck couplings, it is 2 that performs poorly, but only with certain substrates, and the persnickety behavior of this palladium complex is not really understood. Polymer 3 can also pose a challenge when it comes to serving as a precursor to palladacyclic complexes. These are frequently prepared at lower temperature and so solubility could again be the root cause of the issue. Adding to the overall theme, hydrolysis product 4 has been shown to be an issue in Wacker oxidation.123 Given that impurities are so prevalent and can impact the outcome of a reaction involving commercially available palladium acetate, these reports emphasize the importance of knowing the purity of a metal complex before using it and also appropriate storage prior to use. To coin the title of a catalysis-focused article published in 2003, “taking too many precautions in making a catalyst is never a loss of time.”111 The group stated that this was a lesson they learned at their own expense. Until they went back and critically assessed their preparative route, they did not realize that the nuclearity of palladium acetate clusters had a significant impact on the activity of the complexes they prepared and used.

1.21.4.3

Reagent impurities

From their beginnings in organometallic chemistry, metal alkylidene complexes have revolutionized preparative organic chemistry due to their role as alkene metathesis catalysts. They are used in applications as diverse as polymer synthesis, biological chemistry, and drug discovery.125 Of the metal complexes that are currently used as alkene metathesis catalysts, those containing ruthenium predominate.126 This is because they are generally the most rugged; they tolerate, at least to some degree, moisture and air, as well as a variety of functional groups on the alkene substrate. This said, the deleterious effect of impurities in the reaction mixture on the catalytic activity of ruthenium complexes has been known for some time.127 One recent case where impurities in the substrate itself negatively impact the outcome of the reaction involves the ring-opening metathesis polymerization (ROMP) of functionalized norbornene monomers (Fig. 12A).128 The monomers, formed by a copper-catalyzed reaction, themselves originate from an alkynefunctionalized norbornene starting material, 1. Traces of this reagent that remain in the monomer solution prove to be the root cause of the issues in the ROMP process. The alkyne moiety of 1 can react with the ruthenium alkylidene catalyst, 2, to generate a vinyl carbene complex, 3. This complex is a significantly poorer alkene metathesis catalyst as compared to 2, the consequence of which is that the rate of ROMP is significantly retarded. This is actually not unexpected; vinyl carbene complexes are known to be notoriously slow catalysis in enyne metathesis reactions.112 Staying with the theme of ROMP, another example where substrate purity is key is in the preparation of cyclic polybutadienes using 1,5-cyclooctadiene (COD) as the monomer unit.129 Traces of 4-vinylcyclohexene, an impurity left over from the industrial preparation of COD, change the path of the ROMP reaction and lead to linear rather than cyclic products (Fig. 12B). The impact of impurities on the outcome of alkene metathesis reactions is not always negative. In one recent case impurities actually enhance the reaction.113 The ring-closing metathesis (RCM) of diethyl diallylmalonate, 5, can lead to either the desired internal cycloalkene product, 6, or the off-target alkene-functionalized cycloalkane 7 (Fig. 13).114 When using (N-heterocyclic carbene)Ru(arene)Cl2 complexes as catalysts in conjunction with Lewis acids and non-coordinating salts as additives, the purity of the diallylmalonate reagent has a significant impact on the selectivity of the reaction. If ultrapure 5 is used, a significant quantity of unwanted cycloisomer 7 is formed. However, analysis of a range of commercially available batches of diallylmalonate shows that some contain over 100 impurities, many of which are organohalide species. These enhance the metathesis reaction to different degrees and not only is the selectivity for desired product 6 improved, but the overall yield is also increased. The role of the organohalide is to generate an active metathesis catalyst. The diallylmalonate substrate first reacts with the catalyst precursor in an oxidative cyclization process to generate a metallacycle intermediate and then mechanistic bifurcation occurs. In the case of ultrapure diallylmalonate, a b-hydride elimination pathway leads to generation of a ruthenium hydride complex which is a cycloisomerization catalyst and leads to formation of 7. If an organohalide is present, it reacts with the Lewis acid additive present in the mixture to form a carbocation, which abstracts an a-hydride from the metallacycle intermediate, resulting in an alkylidene complex that serves as a metathesis catalyst, which in turn generates 6.130 All these cases certainly show the importance of assaying feedstocks for metathesis reactions prior to use.

646

Impurities in Organometallic Catalysis

Fig. 12 Impact of reagent impurities on ruthenium-catalyzed ring-closing metathesis reactions.

1.21.4.4

Solvent impurities

In any reaction mixture, it is generally the solvent that is present in the greatest amount. As a result, even small impurities in this component could feasibly have the most significant effect on the outcome of the reaction. In many cases, reactions are fairly tolerant of trace solvent impurities, but this is not always the case. One example is in ruthenium-catalyzed direct arylation, one of an evergrowing group of processes involving CdH bond activation.131 This particular atom-economic and streamlined variant of traditional cross-coupling reactions involves the replacement of the organometallic coupling partner with the CdH bond of an arene.132 An aryl halide is the second coupling partner, and a base is used to remove the halogen group from a ruthenium halide intermediate. It is not unusual to see a solvent screen as part of any report of a new metal-catalyzed methodology. Often little emphasis is placed on interpreting the results, except to comment on broad classifications such as polar versus non-polar and protic versus aprotic. Like many reactions before them, ruthenium-catalyzed direct arylations have been performed in a range of different solvents. Generally N-methylpyrrolidinone (NMP) has proven superior but solvents with similar properties, such as dimethylformamide and dimethylacetamide (DMF and DMA) are much less effective, which should have raised suspicion. It was not until a surprising result was observed that a possible rationalization for the efficacy of NMP in this class of reaction was proposed (Fig. 14). In the development of a route to anacetrapib, a drug candidate being developed to treat elevated cholesterol levels, a direct arylation was the linchpin.133 Although the reaction proceeded smoothly when using ruthenium catalyst loadings of ca. 5 mol%, when this was reduced to 0.5–1 mol%, the reaction became highly dependent on the source of the NMP, with product yield ranging from 30% to 90%. Analysis of “good” lots of NMP showed that quantities of g-butyrolactone were present. Given that NMP is generally prepared industrially by treating g-butyrolactone with methylamine, the presence of this contaminant is not surprising.115 In the direct arylation process, the base present in the reaction mixture can facilitate the hydrolysis of g-butyrolactone to generate

Impurities in Organometallic Catalysis

Fig. 13 Enhancement of a ruthenium-catalyzed ring-closing metathesis reaction by reagent impurities.

Fig. 14 Impact of solvent impurities on a ruthenium-catalyzed coupling reaction.

647

648

Impurities in Organometallic Catalysis

carboxylate salts which in turn can enhance the efficiency and stability of the ruthenium catalyst. This is supported by previous anecdotal evidence of the efficacy of NMP in ruthenium-catalyzed direct arylation,134 results showing that addition of carboxylic acids improves the outcome of this class of reaction135,136 and now the use of well-defined ruthenium carboxylate complexes for the transformation.137 This example shows that a solvent cannot be simply considered as a medium for performing a reaction or as an innocent bystander, but instead can bring impurities that influence catalytic processes, be that positively or negatively.

1.21.4.5

The particular case of carbon materials

In many modern applications, carbon-based materials reign supreme; their physical properties outshine those of the alternatives.138 Within this portfolio, graphene and its cousins, carbon nanotubes (CNTs), find use in electronic, optical, biosensing, and biomedical devices.139–141 More relevant to the discussion here, carbon nanotubes have been used extensively in catalysis.142 This has led to a new subfield—carbon-based metal-free catalysis.143,144 However, things may not be as they seem. The traditional routes to CNTs involve the use of metal precursors, particularly iron, nickel, cobalt, and molybdenum.145 These metals serve as catalysts for the chemical vapor deposition (CVD) growth of the nanotubes. Graphene can also be produced this way, or from graphite using permanganate oxidation followed by a reduction/doping step.142,146 As a result, both CNTs and graphene can be contaminated with metal impurities.147 Removal of these poses a significant challenge. The most commonly used approach is to wash the materials with nitric acid which is expected to remove residual metallic components as well as some amorphous carbon but turns out not to be particularly effective.148 The kinetics of dissolution of many metal oxides and carbides, even in concentrated mineral acids, can be very slow.149 As a result, metals are left behind after the washing. These impurities alter the properties of the carbon-based materials quite dramatically (Fig. 15).150,151 This is particularly the case in the area of electrocatalysis.152 Indeed, metal nanoparticles embedded in CNTs may be responsible for most, if not all, of the electrocatalytic activity.153 This is seen in the electroreduction of hydrogen peroxide154–156 and ferricyanide,155 and the oxidation of organic substrates.157 The use of CNTs as electrochemical sensors is also impacted.158 The same issues occur when using graphene.159 In some cases, the exact origin of the electrochemical activity is hard to elucidate. In one intriguing case, the electrochemical activity in the reduction of hydrogen peroxide by double-walled CNTs has been traced back to iron impurities present in the copper catalyst used for CVD fabrication of the materials.153 That is to say, the activity is due to an impurity in the impurity. Together, these examples clearly sound alarm bells and raise significant challenges when determining whether processes involving carbon-based materials truly are “metal-free” and, if they are not, what exactly is responsible for the activity. Since it is central to the application of these materials, this is clearly an important topic for further study. The challenges associated with detecting impurities are discussed later in the article.

Fig. 15 Metal impurities in carbon materials.

Impurities in Organometallic Catalysis

1.21.5

Detecting and avoiding impurities

1.21.5.1

Introduction

649

Given the impact they can have on a reaction, the avoidance, or at the very least the detection of even very low levels of impurities is key. With the spate of examples in the early- to mid-2000s, responsible chemists have justifiably become more skittish when it comes to making claims regarding unusual catalytic activity, and reports of “metal-free” variants of reactions that are traditionally catalyzed by metals have received more scrutiny. This is evidenced by the observations that analysis of substrates, equipment, and entire reaction mixtures for impurities has seen a significant uptick, and that methods for trying to circumvent contamination have come to the fore. Some of the key tenets are highlighted here.

1.21.5.2

Analytical tools

The minimum level of detection is key when it comes to employing analytical techniques as tools for determining whether impurities could be playing a role in a reaction. Of all the techniques available, inductively coupled plasma mass spectroscopy (ICP-MS) is by far the most widely used.160 The ability to detect metals down to the parts per trillion (ppt) range in a reliable and reproducible manner makes it particularly apt to this application. However, it needs to be used appropriately. In many literature reports, single reagents or metal complexes are analyzed. While this is a starting point, it is also necessary to assay entire reaction mixtures. There have been a number of instances of adventitious metal contamination coming from what may seem like otherwise innocuous sources. A case in point is that Heck couplings can be performed without the addition of further catalyst in glassware previously used for the reaction.161 The thiol-ligated palladium complex PdCl2(SEt3)2 can serve as a catalyst precursor for the Heck coupling of iodoarenes with methyl acrylate. If the glassware and stir bar for the reaction are washed, a second dose of reagents and base can be added and the reaction will still be successful, even without any palladium added. Indeed, ICP-MS shows that the reaction can be performed even if ppt levels of the palladium precursor remain in the vessel (Fig. 16). This then begs the question as to how low you should go when it comes to analyzing reaction mixtures. At what point can it be stated definitively that catalytic activity does not arise from impurities in a component of the reaction or adventitious metal present in the apparatus used? There is a general guiding principle that can be applied, based on fundamental kinetics. The rate of a reaction in solution cannot exceed the diffusion limit (109 M−1 s−1 for a bimolecular reaction), and this means there is essentially a limit on catalytic rate. If the amount of impurity is lower than the level that could produce the observed catalytic rate in the reaction within this constraint, then the impurity cannot be responsible. While certainly useful, there are some drawbacks to using ICP-MS as an analytical tool. It requires quite costly apparatus and unlike gas chromatography–mass spectroscopy (GC-MS), or nuclear magnetic resonance (NMR) and ultraviolet-visible (UV-vis) spectroscopy, it is not generally amenable to a hands-on format for the non-expert. Another issue is that full digestion of the sample is needed, which in turn requires that a truly representative sample is placed in the digesting tube. With samples potentially being in a combination of forms, such as a solution containing a slurry or a solid, the technique can sometimes fall foul of the analysis not being wholly representative of the reaction or product mixture. In addition, ICP-MS equipment can have a “memory effect”; one sample can impact the analysis of the next.162 This brings with it the challenge of erroneous results and decreased sensitivity with time. One alternative technique that has seen some interest is neutron activation analysis (NAA).149,163,164 This can be used for detection of metals down to the ppb level. The most significant advantage of NAA is the ability to assay metal content regardless of the format of the sample. It can be used to perform analyzes on solids, solutions, suspensions, and slurries with little or no preparation. Due to the penetrating nature of incident neutrons and resultant gamma rays, the technique provides a true bulk analysis. This said, it too has drawbacks. By its very nature, it requires a neutron source which generally means it needs to be performed in a dedicated facility. In addition, the irradiated sample remains radioactive for many years after the initial analysis, meaning that costly handling and disposal protocols are required.

Fig. 16 Residual palladium in a flask impacting a Heck coupling reaction.

650

Impurities in Organometallic Catalysis

Focusing on carbon-based materials, the detection of metal impurities poses a particular challenge. The metals, coming from the preparation of nanotubes and graphene, are often intercalated deep in the materials. As a result, the traditionally employed techniques for detection such as X-ray photoelectron spectroscopy or energy-dispersive X-ray spectroscopy are not really appropriate.148 They only probe the surface of the materials and, even then, might not be sensitive enough to detect trace levels of metal nanoparticles.148,152 The same is true of another often relied upon technique, thermogravimetric analysis (TGA). Not only is this ineffective, but it also does not offer any element-specific information.165 One technique that does seem to offer valuable information about the level of metal contaminants is direct current (dc) magnetic susceptibility measured with a superconducting quantum interference device (SQUID) magnetometer.166,167 The SQUID technique does not differentiate between bulk and surface properties and measures the total magnetic moment of the sample.168 In a side-by-side comparison of a range of analytical techniques, dc magnetic susceptibility was the only tool that was effective, others providing false “impurity-free” information.167 With the issues of both the cost of and access to some analytical techniques, another approach is to use more readily available tools to probe for trace metals in reaction mixtures. One such example is the application of UV-Vis and fluorescence spectrophotometry.169,170 This proves particularly useful for the identification of palladium and is amenable to a well-plate approach and hence accelerated and efficient screening.171–179 In one approach to these assays, a non-fluorescent reagent is added to a solution of the analyte. A palladium-catalyzed reaction then takes place to generate a fluorescent product which is then detected using spectrophotometry. The palladium-catalyzed CdO bond cleavage of allyl ethers (Tsuji-Trost reaction) has been used in this endeavor, fluorescein-or phenoxazine-derived sensors being employed (Fig. 17A).167–173 The reaction is performed in the presence of a phosphine which not only serves as a ligand but also reduces any Pd(II) to Pd(0) meaning that a total palladium content can be determined, regardless of oxidation state of the metal. By using appropriate conditions, the technique can be employed to determine palladium concentrations down to 0.1 ppb. The fluorescence-based approach can be extended to visualize residual palladium in laboratory glassware, the flask showing a strong green fluorescence when placed above a UV lamp (Fig. 17B).180 Fluorescence assays have been used to probe a purportedly “no catalyst added” Suzuki coupling reaction (Fig. 17C).181 The coupling of aryl iodides with boronic acids was performed using tripotassium phosphate as a base and dimethyl carbonate as the solvent but without an exogenous metal catalyst.182 A fluorescence assay later showed the presence of palladium at a concentration

Fig. 17 A fluorescence-based approach to Probing metal impurities.

Impurities in Organometallic Catalysis

651

of approximately 0.8 ppb. Whether or not this quantity of palladium is enough to catalyze the reaction is still a matter of debate. It is well below the 50 ppb previously reported to be catalytically active, and is certainly an example of “how low should you go?” when it comes to detecting metals such as palladium which could be a catalyst for a reaction. While the use of fluorescence assays has proven useful for interrogating reaction mixtures for residual metal, there are drawbacks that need to be taken into account. Remaining with the theme of palladium, the presence of base in the analyte solution can pose a challenge as can variations in the parameters of the Tsuji-Trost reaction used to generate the fluorescent species in the visualization reaction. Another problem is that the majority of the fluorescence-based approaches to palladium detection do not work effectively when other metals are present in significant amounts. While many of these issues are now better understood and solutions or workarounds have been found to overcome them,168,170,173 results still need to be viewed with caution, and combining complementary techniques for probing metal contaminants is a prudent approach rather than relying solely on one tool.

1.21.5.3

Approaches to avoiding impurities

In light of the number of examples of “impurity-catalyzed” reactions, the rapid rate of publication of “metal-free” reactions that have traditionally been performed using a metal catalyst can be somewhat alarming.183 In the majority of cases, these reactions may truly be metal-free but it is good practice to take precautionary measures and/or perform a series of tests to prove this assertion. There are a variety of options available to chemists (Table 3). One is to use an analytical tool to assay the reaction mixture for metals, as described above. Performing the reaction using new glassware and stir bars is also a worthwhile approach.184,185,189 This obviates some of the potential issues arising from potential contamination coming from previous use of the apparatus in reactions involving metal sources. Washing used glassware is not completely effective, even when using aqua regia (mixture of concentrated hydrochloric and nitric acids), piranha solution (mixture of sulfuric acid, water, and hydrogen peroxide) or a base bath (ethanol or isopropanol and sodium or potassium hydroxide).149,170,171 Stir bars have proven to be particularly pernicious.179 The surface of magnetic stir bars is susceptible to abrasions caused by friction between the stir bar and the inner surface of the glassware, even after just one use. The chance of abrasion is heightened when solids such as insoluble bases and heterogeneous catalysts are used. These abrasions can trap trace quantities of components from a reaction mixture and washing does not effectively remove them. This is a particular issue when it comes to metal species because they are attracted to materials like PTFE that are commonly used as stir bar coatings. As a result, performing control experiments with pristine stir bars is important when testing metal-free or ultralow metal loading protocols, to ensure that the reaction is not “stir bar catalyzed.” One source of contamination can be at least limited if not avoided entirely either by preparing starting materials using a route that does not involve metal-containing reagents or catalysts, or by purifying them before use (Fig. 18). Given the number of examples where bases contain transition-metal impurities, these reagents need particular attention.5,53,71,72,95 As an example, the sublimation of potassium tert-butoxide has been performed as a first line of defense in development of two metal-free CdH bond arylation protocols.186,190,191 In some cases, purification of the base is not trivial. In the synthesis of benzimidazoles by basemediated intramolecular N-arylation of amidines, the starting materials were prepared by a metal-free route and the potassium hydroxide used was of >99.99% purity. Table 3

Approaches to avoiding metal impurities.

Reaction

Tests performed/Steps taken

Conjugation addition of organozinc halides to enones

• • • • • • • • • • • • • • • • • • • • •

Synthesis of Benzimidazoles

a-Arylation of enolizable aryl ketones CdH Bond arylation

Hydroxymethylation of isoquinolines

Cyclization of o-nitroaryl ynamides and ynamines

Boryl substitution

Glassware and stir bar washed with hot aqua regia Trials conducted with new glassware and stir bar Zinc powder assayed for metal impurities Preparation of starting materials without using transition-metals Reagent transfers with plastic spatulas New glassware used for cyclizations Potassium hydroxide of >99.99% purity Resublimation of potassium t-butoxide base Performing reaction with added palladium for comparison Sublimation of potassium t-butoxide and then ICP-MS & ICP-AES Glassware and equipment thoroughly cleaned prior to use Kinetic study finding zero-order dependence on transition-metals Control reaction with new stir bar Use of metal scavenger resin ICP-MS of base and two representative products Use of metal scavenger resin ICP-MS of representative reaction mixture Kinetic studies of reaction with and without trace transition-metals ICP-MS of potassium methoxide base Comparing reaction with and without trace transition-metals Reaction repeated by different research group using new glassware and reagents

References 180

181

182 183,184

185

186

187,188

652

Impurities in Organometallic Catalysis

Fig. 18 Examples of the role of reagent purification in avoiding metal contamination.

Another way to mitigate contamination by transition metals is to use a scavenger reagent. The objective is to remove any metal present in the reaction mixture.192 This has found particular application in the case of palladium, a 2,4, 6-trimercaptotriazine-functionalized silica scavenger (Si-TMT) being used in a number of cases.187,188 If the outcome of performing reactions in the presence of the scavenger is the same as in its absence, this adds additional evidence supporting the assertion that the process is palladium-free. However, since scavengers do not eradicate metals from reaction mixtures completely, this approach should not be used as a stand-alone tool. Another tactic is, counterintuitively, to add metal to the reaction being scrutinized. If performing the reaction in the presence of a plausible impurity catalyst does not change the outcome or kinetics of the transformation, this adds to the credence that the process is indeed not catalyzed by that metal species (Fig. 19). This approach has been taken in a number of instances, particularly in the case of arylation reactions.182,186,193 For example, the a-arylation of enolizable aryl ketones is traditionally performed using a palladium catalyst.194 When a base-mediated transition-metal free variant is performed in the presence of bis(dibenzylideneacetone)palladium, Pd(dba)2, the reaction essentially shuts down.182 This result has been used to suggest that the protocol is a radical-mediated metal-free process. In an analogous methodology where 10 ppb–10 ppm Pd, Cu, Fe and other metals were found in the base and additive, the addition of up to 1000-fold quantities of these same metal salts does not enhance the outcome of the reaction, and in some cases has a deleterious effect.195 One further example of this approach is in the reaction of 1-dibromovinyl-2-nitro-substituted arenes with secondary amines.186 Addition of 10 ppm of a variety of transition metal salts has no impact on the outcome of the reaction and this observation being used to support the working hypothesis that the reaction proceeds by a metal-free cascade mechanism through a key ynamine intermediate. While all these studies may be valid, there is one underlying issue. What if the metal salt added to the reaction mixture is not as active as the metal in the impurity? As a result, other complimentary techniques also need to be used to corroborate the evidence. A final approach that merits mentioning is that of collaboration. As mentioned in the introduction, alarm bells were rung when the “transition-metal metal-free” Suzuki coupling reported by a group in the United Kingdom could not be repeated in Sweden or the United States, the reason being that the sodium carbonate sourced in the UK came laced with parts-per-billion of palladium.5 This was discovered after the protocol had been published. Enlisting the help of another research group to test a protocol before publication helps to validate the results. A case in point is in the development of a transition-metal free base-mediated methodology for boryl substitution using silylboranes.196,197 Control experiments with new reagents and apparatus conducted by another research group showed almost the same results.

Impurities in Organometallic Catalysis

653

Fig. 19 Probing the role of impurities on a reaction by adding metals.

1.21.6

Concluding remarks

This article has shown that sometimes catalytic reactions are not what they seem at first glance; the possible role that impurities can play in catalysis must not be overlooked. A number of examples have been showcased here. In hindsight, some may seem obvious, but the appropriate skepticism can be easily forgotten in the excitement of finding a “new” catalyst. Though discovery of catalysis effected by impurities may seem to be disastrous, such discoveries have also opened new avenues of research. An example from the field of medicinal chemistry is the discovery that platinum salts can have potent anti-cancer activity.198,199 This discovery started with a hypothesis about the influence of electric fields on organisms, and an experiment showed that cells stopped dividing when subjected to an electric field. The effect was later found to be solely due to trace quantities of platinum(II) from the platinum electrodes that dissolved in the electrolysis medium and reacted to form cis-[PtCl2(NH3)2].200 Later this complex, known as cisplatin, was also found to block cell division in tumors. It and later variants are now used widely as chemotherapeutic agents and have saved numerous lives.201,202 Returning to catalysis, the door has also opened to operationally simpler methodologies. For example, the fact that the Suzuki coupling can be performed with sub-ppm levels of palladium catalysts means that costly metal removal procedures can potentially be avoided at the end of the reaction.5 In the pharmaceutical industry, there is a maximum threshold of 10 ppm for residual palladium in a final product.203 If the catalyst loading is below this level to begin with, purification of the product could feasibly be simplified. Alongside this, more sustainable approaches can be applied to synthetic chemistry; low catalyst loadings by their very nature mean that less of the precious metal is used. Operating at these ultra-low catalyst loadings is beneficial in terms of pushing catalyst efficiency to the maximum. For some reactions catalyzed by ligand-free palladium sources, the formation of inactive metal aggregates is an issue.46,92,204 When working at very low catalyst concentrations, this deactivation route can be avoided, or at least ameliorated.

654

Impurities in Organometallic Catalysis

In some cases, catalytic activity is due to an impurity hitching a ride, as in the numerous reports of purportedly “iron-catalyzed” coupling reactions that upon reinvestigation turned out to be copper or palladium mediated.3,27,56,57 Another example where impurities driving catalysis is prevalent is in the case of the poster child of modern materials: carbon-based materials, specifically nanotubes and graphene.148–153 However, impurities need not necessarily be wholly responsible for observed catalytic activity. Sometimes they enhance the activity of another catalyst; as described above, carboxylate salts formed by the hydrolysis of g-butyrolactone, an impurity in N-methyl pyrrolidone, enhance ruthenium-catalyzed alkene metathesis.115 They can also change the outcome of a reaction. In the ring-opening metathesis polymerization of cyclooctadiene, traces of an impurity left over from the industrial preparation of this feedstock changes the path of the reaction, leading to linear rather than cyclic products.113 The increasing number of examples of impurity-catalyzed and impurity-modified reactions has led to more rigorous assessment of methodologies. This takes a number of different guises. Chemists are using analytical techniques such as ICP-MS and fluorescence spectrophotometry to probe for rogue elements.160,165–177 They are using new glassware and stir bars,179 purifying and then repurifying reagents or preparing them in new ways,181–184 adding metals to see if they impact the reaction,182–184,186–188 or turning to collaborators to prove reproducibility,187,188 all with the aim of probing whether “metal-free” is really metal-free. New reports of methodologies are being more thoroughly assessed in the review stage prior to publication. This scrutiny leads to robust, reliable reaction chemistry. While this is the first time a article on the topic of “metal impurities in catalysis” has appeared in Comprehensive Organometallic Chemistry, there will undoubtably be future examples that will fill the pages of further reviews. Indeed, in the process of writing this current article, the specter of “transition-metal free” Suzuki coupling has raised its head again. In early 2021 a report emerged of a Suzuki-type coupling of aryl halides with arylboronic acids catalyzed by amines.205 The coupling appeared to work well for a range of substrates and a putative mechanism was proposed based on experimental and computational analysis. A range of the standard tools were used as due diligence to test for the presence of palladium in the reaction. Given all the literature precedent, and the advent of social media as a means of exchanging ideas, almost immediately after the appearance of the work on the journal website, people started to debate whether the reaction was truly “palladium-free.” Some supported the authors and others were adamant that trace palladium was at play. Within weeks, another modern tool for rapid dissemination of results, an open access preprint archive for chemistry, came into play. An article on the site reported results suggesting that palladium was introduced as a contaminant left over from Buchwald-Hartwig couplings that were used to make the amine “catalysts” for the Suzuki-type reaction.206 Then a second preprint appeared corroborating this.207 An international team of academic and industrial chemists conducted a reinvestigation of the key claims in the original article and concluded that the observed catalytic activity could not be due to the amine. Instead, tricyclohexylphosphine palladium complexes that are readily entrained during the purification of the amine were the culprit. With all this said, while the impact of impurities in the catalysis field is certainly a call for carefulness, it is important that it does not dampen creativity and innovation. In many cases “metal-free” is really “metal free” and non-traditional metals really do catalyze reactions.178,208 A later discovery that an impurity is the culprit, although not ideal, need not necessarily be the end of the road and invalidate the work entirely. It might instead be the stepping-stone to something even more interesting.

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.

Heim, M. Exploring Indiana Highways: Trip Trivia; Travel Organization Network Exchange Inc.: Wabasha, Minnesota, USA, 2007; p 66. O’Brien, J.; Lehtonen, K. Counterfeit Mobile Devices – The Duck Test. In 10th International Conference on Malicious and Unwanted Software (MALWARE), 2015; pp 144–151. Thomé, I.; Nijs, A.; Bolm, C. Chem. Soc. Rev. 2012, 41, 979–987. Leadbeater, N. E. Nat. Chem. 2010, 2, 1007–1009. Arvela, R. K.; Leadbeater, N. E.; Sangi, M. S.; Williams, V. A.; Granados, P.; Singer, R. D. J. Org. Chem. 2005, 70, 161–168. Connell, T. U.; Schieber, C.; Proietti Silvestri, I.; White, J. M.; Williams, S. J.; Donnelly, P. S. Inorg. Chem. 2014, 53, 6503–6511. For a contemporaneous review, see: Kharasch, M. S.; Reinmuth, O. Grignard Reactions of Nonmetallic Substances, Prentice-Hall, New York, USA, 1954. Cusa, N. W.; Kipping, F. S. J. Soc. Chem. Ind. Lond. 1934, 53, 213–214. Kharasch, M. S.; Fields, E. K. J. Am. Chem. Soc. 1941, 63, 2316–2320. Kharasch, M. S.; Tawney, P. O. J. Am. Chem. Soc. 1941, 63, 2308–2316. For a historical perspective, see: Claverie, J.; Schaper, F. MRS Bull. 2013, 38, 213–218 Ziegler, K. Angew. Chem. 1952, 64, 323–329. For an overview, see: Fischer, K.; Jonas, K.; Misbach, P.; Stabba, R.; Wilke, G. Angew. Chem. 1973, 12, 943–953. (a) Ziegler, K.; Holzkamp, E.; Breil, H.; Martin, H. Angew. Chem. 1955, 67, 426; (b) Ziegler, K.; Holzkamp, E.; Breil, H.; Martin, H. Angew. Chem. 1955, 67, 541–547. Jin, H.; Uenishi, J. U.; Christ, W. J.; Kishi, Y. J. Am. Chem. Soc. 1986, 108, 5644–5646. Takai, K.; Tagashira, M.; Kuroda, T.; Oshima, K.; Utimoto, K.; Nozaki, H. J. Am. Chem. Soc. 1986, 108, 6048–6050. For a personal account, see: Takai, K. Bull. Chem. Soc. Jpn. 2015, 88, 1511–1529. Takai, K.; Kakiuchi, T.; Utimoto, K. J. Org. Chem. 1994, 59, 2671–2673. Jubert, C.; Knochel, P. J. Org. Chem. 1992, 57, 5431–5447. Okazoe, T.; Takai, K.; Oshima, K.; Utimoto, K. J. Org. Chem. 1987, 52, 4410–4412. Takai, K.; Kakiuchi, T.; Kataoka, Y.; Utimoto, K. J. Org. Chem. 1994, 59, 2668–2670. Leadbeater, N. E.; Marco, M. Angew. Chem. Int. Ed. 2003, 42, 1407–1409. Leadbeater, N. E.; Marco, M. J. Org. Chem. 2003, 68, 5660–5667. Kylmälä, T.; Valkonen, A.; Rissanen, K.; Xu, Y.; Franzén, R. Tetrahedron Lett. 2008, 49, 6679–6681. Bedford, R. B.; Nakamura, M.; Gower, N. J.; Haddow, M. F.; Hall, M. A.; Huwe, M.; Hashimoto, T.; Okopie, R. A. Tetrahedron Lett. 2009, 50, 6110–6111. Kylmälä, T.; Valkonen, A.; Rissanen, K.; Xu, Y.; Franzén, R. Tetrahedron Lett. 2009, 50, 5692. Buchwald, S. L.; Bolm, C. Angew. Chem. Int. Ed. 2009, 48, 5586–5587. Gonda, Z.; Tolnai, G. L.; Novák, Z. Chem. A Eur. J. 2010, 16, 11822–11826.

Impurities in Organometallic Catalysis 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. 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.

655

Zuidema, E.; Bolm, C. Chem. A Eur. J. 2010, 16, 4181–4185. Plenio, H. Angew. Chem. Int. Ed. 2008, 47, 6954–6956. González-Arellano, C.; Abad, A.; Corma, A.; García, H.; Iglesias, M.; Sánchez, F. Angew. Chem. Int. Ed. 2007, 46, 1536–1538. Lauterbach, T.; Livendahl, M.; Rosellón, A.; Espinet, P.; Echavarren, A. M. Org. Lett. 2010, 12, 3006–3009. Corma, A.; Juárez, R.; Boronat, M.; Sánchez, F.; Iglesias, M.; García, H. Chem. Commun. 2011, 47, 1446–1448. For reviews, see: (a) Nijamudheen, A.; Datta, A. Chem. A Eur. J. 2020, 26, 1442–1487.(b)(b) Kramer, S. Synthesis 2020, 52, 2017-2030.(c) Shi, Q.; Qin, Z.; Xu, H.; Li, G. Nanomaterials 2019, 9, 838; (d) Akram, M. O.; Banerjee, S.; Saswade, S. S.; Bedi, V.; Patil, N. T. Chem. Commun. 2018, 54, 11069–11083; (e) Joost, M.; Amgoune, A.; Bourissou, D. Angew. Chem. Int. Ed. 2015, 54, 15022–15045; (f ) Stratakis, M.; Garcia, H. Chem. Rev. 2012, 112, 4469–4506; (g) Zhang, Y.; Cui, X.; Shi, F.; Deng, Y. Chem. Rev. 2012, 112, 2467–2505; (h) Wegner, H. A.; Auzias, M. Angew. Chem. Int. Ed. 2011, 50, 8236–8247 Santilli, C.; Beigbaghlou, S. S.; Ahlburg, A.; Antonacci, G.; Fristrup, P.; Norrby, P.-O.; Madsen, R. Eur. J. Org. Chem. 2017, 5269–5274. Kang, S.-K.; Kim, J.-S.; Choi, S.-C. J. Org. Chem. 1997, 62, 4208–4209. (a) Ayouchia, H. B. E.; Bahsis, L.; Fichtali, I.; Domingo, L. R.; Ríos-Gutiérrez, M.; Julve, M.; Stiriba, S.-E. Catalysts 2020, 10, 956; (b) Sasidharan, D.; Namitha, T. R.; Johnson, S. P.; Jose, V.; Mathew, P. Sustain. Chem. Pharm. 2020, 16, 100255. For reviews, see: (a) Sanjosé-Orduna, J.; Mudarra, Á. L.; Martínez de Salinas, S.; Pérez-Temprano, M. H. ChemSusChem 2019, 12, 2882–2897.(b) Cheng, G.-J.; Zhong, X.-M.; Wu, Y.-D.; Zhang, X. Chem. Commun. 2019, 55, 12749–12764; (c) Sperger, T.; Sanhueza, I. A.; Schoenebeck, F. Acc. Chem. Res. 2016, 49, 1311–1319; (d) Bonney, K. J.; Schoenebeck, F. Chem. Soc. Rev. 2014, 43, 6609–6638 For perspective, see: (a) de Meijere, A.; Bräse, S.; Oestreich, M. Eds. Metal-Catalyzed Cross-Coupling Reactions and More, Wiley-VCH, Weinheim, Germany, 2013; (b) Vrijdag, J., Ed.; In Cross-Coupling Reactions: An Overview; Nova Science Publishers: Hauppauge, NY, 2020; (c) Colacot, T., Ed.; In New Trends in Cross-Coupling: Theory and Applications; Royal Society of Chemistry: Cambridge, UK, 2014; (d) Kostas, I. D., Ed.; In Suzuki-Miyaura Cross-Coupling Reaction and Potential Applications; MDPI: Basel, Switzerland, 2017. Campeau, L.-C.; Hazari, N. Organometallics 2019, 38, 3–35. For select reviews, see: (a) Chatterjee, A.; Ward, T. R. Catal. Lett. 2016, 146, 820–840.(b) Hoffmann, I.; Blumenröder, B.; Onodi, S.; Dommera, S.; Schatz, J. Green Chem. 2015, 17, 3844–3857; (c) Sherwood, J.; Clark, J. H.; Fairlamb, I. J. S.; Slattery, J. M. Green Chem. 2019, 21, 2164–2213 For a review, see: Jose, D. E.; Kanchana, U. S.; Mathew, T. V.; Anilkumar, G. J. Organomet. Chem. 2020, 927, 121538. (a) Alimardanov, A.; Schmieder-van de Vondervoort, L.; de Vries, A. H. M.; de Vries, J. G. Adv. Synth. Catal. 2004, 346, 1812–1817; (b) Bedford, R. B.; Blake, M. E.; Butts, C. P.; Holder, D. Chem. Commun. 2003, 466–467. For a review of the application of microwave heating as a tool for palladium-catalyzed cross couplings, see: Salih, K. S. M.; Baqi, Y. Catalysts 2020, 10, 4 For perspective on the use of water as a solvent in conjunction with microwave heating, see: Dallinger, D.; Kappe, C. O. Chem. Rev. 2007, 107, 6, 2563–2591. Leadbeater, N. E.; Marco, M. J. Org. Chem. 2003, 68 (3), 888–892. For a review on the use of water as a solvent for Suzuki coupling reactions, see: Leadbeater, N. E. Chem. Commun. 2005, 2881–2902. For a review of palladium-catalyzed cross-coupling reactions performed at ppm to ppb catalyst loadings, see: Roy, D.; Uozumi, Y. Adv. Synth. Catal. 2017, 360, 602–625. For a review, see: Singer, R. A.; Monfette, S.; Bernhardson, D. J.; Tcyrulnikov, S.; Hansen, E. C. Org. Process Res. Dev. 2020, 24, 909–915. For reviews, see: (a) Sandl S., von Wangelin, A. J. Angew. Chem. Int. Ed. 2020, 59, 5434–5437; (b) Guðmundsson, A.; Bäckvall, J.-E. Molecules 2020, 25, 1349; (c) Piontek, A.; Bisz, E.; Szostak, M. Angew. Chem. Int. Ed. 2018, 57, 1116–11128; (d) Mako, T. L.; Byers, J. A. Inorg. Chem. Front. 2016, 3, 766–790; (e) Guérinot, A.; Cossy, J. Top. Curr. Chem. 2016, 374, 49; (f ) Bauer, E. B. Top. Organomet. Chem. 2015, 50, 1–15 Sheloumov, A. M.; Tundo, P.; Dolgushin, F. M.; Koridze, A. A. Eur. J. Inorg. Chem. 2008, 572–576. Bézier, D.; Darcel, C. Adv. Synth. Catal. 2009, 351, 1732–1736. Bézier, D.; Darcel, C. Adv. Synth. Catal. 2010, 352, 1081. Kumar, L. M.; Ansari, R. M.; Bhat, B. R. Appl. Organomet. Chem. 2018, 32, e4054. (a) Kumar, L. M.; Mishra, P.; Bhat, B. R. Catal. Lett. 2019, 149, 1118–1124; (b) Ansari, R. M.; Bhat, B. R. J. Chem. Sci. 2017, 129, 1483–1490. Tailor, S. B.; Bedford, R. B. Catal. Lett. 2020, 150, 963–968. Tailor, S. B.; Manzotti, M.; Asghar, S.; Rowsell, B. J. S.; Luckham, S. L. J.; Sparkes, H. A.; Bedford, R. B. Organometallics 2019, 38, 1770–1777. Kumar, L. M.; Ansari, R. M.; Bhat, B. R. Appl. Organomet. Chem. 2020, 34, e5671. For select examples, see: (a) Crockett, M. P.; Wong, A. S.; Li, B.; Byers, J. A. Angew. Chem. Int. Ed. 2020, 59, 5392–5397; (b) O’Brien, H. M.; Manzotti, M.; Abrams, R. D.; Elorriaga, D.; Sparkes, H. A.; Davis, S. A.; Bedford, R. B. Nat. Catal. 2018, 1, 429–437; (c) Crockett, M. P.; Tyrol, C. C.; Wong, A. S.; Li, B.; Byers, J. A. Org. Lett. 2018, 20, 5233–5237 For reviews, see: (a) Boit, T. B.; Bulger, A. S.; Dander, J. E.; Garg, N. K. ACS Catal. 2020, 10, 12109–12126; (b) Bisz, E.; Szostak, M. ChemSusChem 2017, 10, 3964–3981 Bistri, O.; Correa, A.; Bolm, C. Angew. Chem. Int. Ed. 2008, 47, 586–588. Correa, A.; Carril, M.; Bolm, C. Angew. Chem. Int. Ed. 2008, 47, 2880–2883. Correa, A.; Bolm, C. Adv. Synth. Catal. 2008, 350, 391–394. Correa, A.; Carril, M.; Bolm, C. Chem. A Eur. J. 2008, 14, 10919–10922. Correa, A.; Bolm, C. Angew. Chem. Int. Ed. 2007, 46, 8862–8865. Bonnamour, J.; Bolm, C. Org. Lett. 2008, 10, 2665–2667. Correa, A.; Elmore, S.; Bolm, C. Chem. A Eur. J. 2008, 14, 3527–3529. Larsson, P. F.; Correa, A.; Carril, M.; Norrby, P. O.; Bolm, C. Angew. Chem. Int. Ed. 2009, 48, 5691–5693. For reviews, see: (a) Kanwal, I.; Mujahid, A.; Rasool, N.; Rizwan, K.; Malik, A.; Ahmad, G.; Shah, S. A. A.; Rashid, U.; Nasir, N. M. Catalysts 2020, 10, 443; (b) Chinchilla, R.; Nájera, C. Chem. Soc. Rev. 2011, 40, 5084–5121; (c) Chinchilla, R.; Nájera, C. In Modern Alkyne Chemistry: Catalytic and Atom-Economic Transformations; Trost, B. M., Li, C. J., Eds.; Wiley-VCH: Weinheim, Germany, 2014; pp 269–298; (d) Chinchilla, R.; Nájera, C. Chem. Rev. 2007, 107, 874–922; (e) Doucet, H.; Hierso, J.-C. Angew. Chem. Int. Ed. 2007, 46, 834–871 For reviews, see: (a) Murashkina, A. V.; Mitrofanov, A. Y.; Beletskaya, I. P. Russ. J. Org. Chem. 2019, 55, 1445–1458; (b) Maaliki, C.; Thiery, E.; Thibonnet, J. Eur. J. Org. Chem. 2017, 209–228; (c) Monnier, F.; Taillefer, M. Angew. Chem. Int. Ed. 2008, 47, 3096–3099 Leadbeater, N. E.; Marco, M.; Tominack, B. J. Org. Lett. 2003, 5, 3919–3922. Appukkuttan, P.; Dehaen, W.; Van der Eycken, E. Eur. J. Org. Chem. 2003, 4713–4716. (a) Tian, W.-F.; He, K.-H.; Li, N.; Liu, F.; Mai, X.; Feng, L.-H.; He, Y.-Q. ChemistrySelect 2020, 5, 4496–4499; (b) Luque, R.; Macquarrie, D. J. Org. Biomol. Chem. 2009, 7, 1627–1632. Carril, M.; Correa, A.; Bolm, C. Angew. Chem. Int. Ed. 2008, 47, 4862–4865. Pan, C.; Luo, F.; Wang, W.; Ye, Z.; Liu, M. J. Chem. Res. 2009, 8, 478–481. (a) Hajipour, A. R.; Abolfathi, P.; Tavangar-Rizi, Z. Appl. Organomet. Chem. 2018, e4353; (b) Petuker, A.; El-Tokhey, M.; Reback, M. L.; Mallick, B.; Apfel, U.-P. ChemistrySelect 2016, 1, 2717–2721; (c) Sindhu, K. S.; Thankachan, A. P.; Thomas, A. M.; Anilkumar, G. ChemistrySelect 2016, 1, 556–559; (d) Tran, N. T.; Cho, C. S.; Sohn, H.-S.; Shim, S. C. Bull. Kor. Chem. Soc. 2011, 32, 1080–1082. (a) Hajipour, A. R.; Khorsandi, Z. Appl. Organomet. Chem. 2020, e5398; (b) Hajipour, A. R.; Khorsandi, Z.; Abeshtian, Z. Inorg. Chem. Commun. 2019, 107, , 107470; (c) Feng, L.; Liu, F.; Sun, P.; Bao, J. Synlett 2008, 1415–1417. Park, S.; Kim, M.; Dong, H. K.; Chang, S. Adv. Synth. Catal. 2004, 346, 1638–1640.

656

Impurities in Organometallic Catalysis

79. (a) Wang, Z.; Zheng, T.; Sun, H.; Li, X.; Fuhr, O.; Fenske, D. New J. Chem. 2018, 42, 11465–11470; (b) Hussain, N.; Gogoi, P.; Khare, P.; Das, M. R. RSC Adv. 2015, 5, 103105–103115; (c) Moghaddam, F. M.; Tavakoli, G.; Rezvani, H. R. Cat. Com. 2015, 60, 82–87; (d) Yi, J.; Lu, X.; Sun, Y.-Y.; Xiao, B.; Liu, L. Angew. Chem. Int. Ed. 2013, 52, 12409–12413; (e) Gallego, D.; Brück, A.; Irran, E.; Meier, F.; Kaupp, M.; Driess, M.; Hartwig, J. F. J. Am. Chem. Soc. 2013, 135, 15617–15626; (f ) Bakherad, M.; Keivanloo, A.; Mihanparast, S. Synth. Commun. 2010, 40, 179–185; (g) Gu, S.; Chen, W. Organometallics 2009, 28, 909–914; (h) Wang, M.; Li, P.; Wang, L. Synth. Commun. 2004, 34, 2803–2812; (i) Yan, J.; Wang, Z.; Wang, L. J. Chem. Res. 2004, 71–73; (j) Wang, L.; Li, P.; Zhang, Y. Chem. Commun. 2004, 4, 514–515; (k) Beletskaya, I. P.; Latyshev, G. V.; Tsvetkov, A. V.; Lukashev, N. V. Tetrahedron Lett. 2003, 44, 5011–5013. 80. (a) Cheval, N. P.; Hoffmann, B.; Dikova, A.; Sirindil, F.; Bertus, P.; Blanc, A.; Weibel, J.-M.; Pale, P. Tetrahedron 2018, 74, 7111–7119; (b) Naikwade, A.; Bansode, P.; Rashinkar, G. J. Organomet. Chem. 2018, 866, 112–122; (c) Sanchez-Sanchez, C.; Orozco, N.; Holgado, J. P.; Beaumont, S. K.; Kyriakou, G.; Watson, D. J.; GonzalezElipe, A. R.; Feria, L.; Fernández Sanz, J.; Lambert, R. M. J. Am. Chem. Soc. 2015, 137, 940–947; (d) Li, P.; Wang, L. Synlett 2006, 2261–2265. 81. (a) Sisodiya, S.; Wallenberg, L. R.; Lewin, E.; Wendt, O. F. Appl. Catal., A 2015, 503, 69–76; (b) Lin, J.; Abroshan, H.; Liu, C.; Zhu, M.; Li, G.; Haruta, M. J. Catal. 2015, 330, 354–361; (c) Li, G.; Jiang, D.-E.; Liu, C.; Yu, C.; Jin, R. J. Catal. 2013, 306, 177–183; (d) Kyriakou, G.; Beaumont, S. K.; Humphrey, S. M.; Antonetti, C.; Lambert, R. M. ChemCatChem 2010, 2, 1444–1449; (e) De Souza, R. O. M. A.; Bittar, M. S.; Mendes, L. V. P.; Da Silva, C. M. F.; Da Silva, V. T.; Antunes, O. A. C. Synlett 2008, 1777–1780. 82. (a) Li, P.; Wang, L.; Wang, M.; You, F. Eur. J. Org. Chem. 2008, 5946–5951; (b) González-Arellano, C.; Corma, A.; Iglesias, M.; Sánchez, F. Eur. J. Inorg. Chem. 2008, 1107–1115; (c) Corma, A.; González-Arellano, C.; Iglesias, M.; Pérez-Ferreras, S.; Sánchez, F. Synlett 2007, 1771–1774. 83. Borah, H. N.; Prajapati, D.; Boruah, R. C. Synlett 2005, 2823–2825. 84. For a review, see: Garcia, P.; Malacria, M.; Aubert, C.; Gandon, V.; Fensterbank, L. ChemCatChem 2010, 2, 493–497. 85. Kanuru, V. K.; Kyriakou, G.; Beaumont, S. K.; Papageorgiou, A. C.; Watson, D. J.; Lambert, R. M. J. Am. Chem. Soc. 2010, 132, 8081–8086. 86. Beaumont, S. K.; Kyriakou, G.; Lambert, R. M. J. Am. Chem. Soc. 2010, 132, 12246–12248. 87. Livendahl, M.; Espinet, P.; Echavarren, A. M. Platin. Met. Rev. 2011, 55, 212–214. 88. For pertinent examples, see: (a) Johansson, N.; Sisodiya, S.; Shayesteh, P.; Chaudhary, S.; Andersen, J. N.; Knudsen, J.; Wendt, O. F; Schnadt, J. J. Phys. Condens. Matter. 2017, 29, 444005; (b) Robinson, P. S. D.; Khairallah, G. N.; Da Silva, G.; Lioe, H.; O’Hair, R. A. J. Angew. Chem. Int. Ed. 2012, 51, 3812–3817; (c) Boronat, M.; Combita, D.; Concepción, P.; Laursen, S.; De Dios López-Castro, J. J. Phys. Chem. C 2012, 116, 24855–24867; (d) Hopkinson, M. N.; Gee, A. D.; Gouverneur, V. Chem. A Eur. J. 2011, 17, 8248–8262 89. For reviews, see: (a) Kurandina, D.; Chuentragool, P.; Gevorgyan, V. Synthesis, 2019, 51, 985–1005.(b) Jagtap, S. Catalysts 2017, 7, 267; (c) Beletskaya, I. P.; Cheprakov, A. V. In New Trends in Cross-Coupling: Theory and Applications; Colocot, T., Ed.; Royal Society of Chemistry: Cambridge, UK, 2015, pp 355–478. ch 9 90. For a contemporaneous review, see: Reetz, M. T.; de Vries, J. G. Chem. Commun. 2004, 1559–1563. 91. de Vries, A. H. M.; Mulders, J. M. C. A.; Mommers, J. H. M.; Henderickx, H. J. W.; de Vries, J. G. Org. Lett. 2003, 5, 3285–3288. 92. (a) de Vries, A. H. M.; Parlevliet, F. J.; Schmieder-van de Vonder-voort, L.; Mommers, J. H. M.; Henderickx, H. J. W.; Walet, M. A. M.; de Vries, J. G. Adv. Synth. Catal. 2002, 344, 996–1002; (b) Reetz, M. T.; Westermann, E.; Lohmer, R.; Lohmer, G. Tetrahedron Lett. 1998, 39, 8449–8452. 93. For a review, see: Christoffel, F.; Ward, T. R. Catal. Lett. 2018, 148, 489–511. 94. (a) Zhang, R.; Sato, O.; Zhao, F.; Sato, M.; Ikushima, Y. Chem. A Eur. J. 2004, 10, 1501–1506; (b) Zhang, R.; Zhao, F.; Sato, M.; Ikushima, Y. Chem. Commun. 2003, 3, 1548–1549. 95. Arvela, R. K.; Leadbeater, N. E. J. Org. Chem. 2005, 70, 1786–1790. 96. Sruthi, P. R.; Anjali, S.; Varghese, N.; Anas, S. J. Organomet. Chem. 2020, 921, 121354. 97. (a) Iyer, S.; Thakur, V. V. J. Mol. Catal. A 2000, 157, 275–278; (b) Hajipour, A. R.; Abolfathi, P.; Tavangar-Rizi, Z. Appl. Organomet. Chem. 2018, 6, 4353. 98. (a) Sobhani, S.; Hosseini Moghadam, H.; Skibsted, J.; Sansano, J. M. Green Chem. 2020, 22, 1353–1365; (b) Hajipour, A. R.; Khorsandi, Z.; Abeshtiani, Z.; Zakeri, S. J. Inorg. Organomet. Polym. Mater. 2020, 30, 2163–2171; (c) Hajipour, A. R.; Rezaei, F.; Khorsandi, Z. Green Chem. 2017, 19, 1353–1361; (d) Zhu, Z.; Ma, J.; Xu, L.; Xu, L.; Li, H.; Li, H. ACS Catal. 2012, 2, 2119–2125; (e) Akhlaghinia, B.; Mohammadinezhad, A. Green Chem. 2017, 19, 5625–5641; (f ) Shao, M.; Peng, K.; Chen, J.; Li, H.; Wang, X.; Zhang, W.; Qi, H. Cat. Com. 2009, 10, 1178–1183. 99. For perspective on manganese catalysis, see: (a) Sortais, J.-B. Manganese Catalysis in Organic Synthesis, Wiley-VCH, Weinheim, Germany, 2020.(b) Carney, J. R.; Dillon, B. R.; Thomas, S. P. Eur. J. Org. Chem. 2016, 3912–3929; (c) Valyaev, D. A.; Lavigne, G.; Lugan, N. Coord. Chem. Rev. 2016, 308, 191–235 100. (a) Cahiez, G.; Gager, O.; Lecomte, F. Org. Lett. 2008, 10, 5255–5256; (b) Rueping, M.; Ieawsuwan, W. Synlett 2007, 247–250; (c) Cahiez, G.; Lepifre, F.; Ramiandrasoa, P. Synthesis 1999, 2138–2144; (d) Alami, M.; Ramiandrasoa, P.; Cahiez, G. Synlett 1998, 325–327. 101. Antonacci, G.; Ahlburg, A.; Fristrup, P.; Norrby, P.-O.; Madsen, R. Eur. J. Org. Chem. 2017, 4758–4764. 102. (a) Yong, F.-F.; Teo, Y.-C. Synlett 2012, 2106–2110; (b) Yong, F.-F.; Teo, Y.-C. Tetrahedron Lett. 2010, 51, 3910–3912; (c) Teo, Y.-C.; Yong, F.-F.; Poh, C.-Y.; Yan, Y.-K.; Chua, G.-L. Chem. Commun. 2009, 6258–6260. 103. For background, see: (a) Chen, Y.; Tong, Z.-R. Eds. Click Chemistry: Approaches, Applications and Challenges, Nova Science Publishers, Hauppauge, NY, 2017.(b) Chandrasekaran, S., Ed.; In Click Reactions in Organic Synthesis; Wiley-VCH: Weinheim, Germany, 2016; (c) Rutjes, F., Fokin, V. V., Eds.; In Click Chemistry: In Chemistry, Biology and Macromolecular Science; Wiley-VCH: Weinheim, Germany, 2009. 104. For an overview, see: (a) Meldal, M.; Diness, F. Trends Chem. 2020, 2, 569–584.(b) Nebra, N.; García-Álvarez, J. Molecules 2020, 25, 2015; (c) Liang, L.; Astruc, D. Coord. Chem. Rev. 2011, 255, 2933–2945; (d) Díez-González, S. Cat. Sci. Technol. 2011, 1, 166–178 105. For an review, see: Gomes, R. S.; Jardim, G. A. M.; de Carvalho, R. L.; Araujo, M. H.; da Silva Júnior, E. N. Tetrahedron 2019, 75, 3697–3712. 106. For an review, see: Sultana, J.; Sarma, D. Catal. Rev. Sci. Eng. 2020, 62, 96–117. 107. For recent examples, see: (a) Garg, A.; Khupse, N.; Bordoloi, A.; Sarma, D. New J. Chem. 2019, 43, 19331–19337; (b) Banerji, B.; Chandrasekhar, K.; Killi, S. K.; Pramanik, S. K.; Uttam, P.; Sen, S.; Maiti, N. C. R. Soc. Open Sci. 2016, 3, 160090; (c) Salam, N.; Sinha, A.; Roy, A. S.; Mondal, P.; Jana, N. R.; Islam, S. M. RSC Adv. 2014, 4, 10001–10012; (d) McNulty, J.; Keskar, K. Eur. J. Org. Chem. 2012, 5462–5470 108. For reviews, see: (a) Liu, X.; Li, B.; Liu, Q. Synthesis 2019, 51, 1293–1310; (b) Lee, Y.-C.; Kumar, K.; Waldmann, H. Angew. Chem. Int. Ed. 2018, 57, 5212–5226; (c) Pellissier, H. Adv. Synth. Catal. 2018, 360, 1551–1583 109. (a) Heller, S. T.; Kiho, T.; Narayan, A. R. H.; Sarpong, R. Angew. Chem. Int. Ed. 2013, 52, 11129–11133; (b) Narayan, A. R. H.; Sarpong, R. Green Chem. 2010, 12, 1556–1559. 110. Wilkerson-Hill, S. M.; Yu, D.; Painter, P. P.; Fisher, E. L.; Tantillo, D. J.; Sarpong, R.; Hein, J. E. J. Am. Chem. Soc. 2017, 139, 10569–10577. 111. Bianchini, C.; Meli, A.; Oberhauser, W. Organometallics 2003, 22, 4281–4285. 112. (a) Fischmeister, C.; Bruneau, C. Beilstein J. Org. Chem. 2011, 7, 156–166; (b) Diver, S. T. Coord. Chem. Rev. 2007, 251, 671–701. 113. Lübbe, C.; Dumrath, A.; Neumann, H.; Schäffer, M.; Zimmermann, R.; Beller, M.; Kadyrov, R. ChemCatChem 2014, 6, 684–688. 114. Lübbe, C.; Dumrath, A.; Neumann, H.; Beller, M.; Kadyrov, R. ChemCatChem 2014, 6, 105–108. 115. (a) Harreus, A. L.; Backes, R.; Eichler, J.-O.; Feuerhake, R.; Jäkel, C.; Mahn, U.; Pinkos, R.; Vogelsang, R. Ullmann’s Encyclopedia of Industrial Chemistry; Wiley: New York, 2011; (b) Ohlbach, F.; Melder, J.-P.; Ross, K.-H.; Rudloff, M.; Liebe, J. US Patent 6348601, 2002; (c) Bertola, A. US Patent 6248902, 2001. 116. Kaufhold, S.; Petermann, L.; Sorsche, D.; Rau, S. Chem. A Eur. J. 2017, 23, 2271–2274. 117. Gallo, V.; Mastrorilli, P.; Nobile, C. F.; Romanazzi, G.; Suranna, G. P. J. Chem. Soc. Dalton Trans. 2002, 4339–4342. 118. Grennberg, H.; Foot, J. S.; Banwell, M. G.; Roman, D. S. Encyclopedia of Reagents for Organic Synthesis; Wiley: New York, USA, 2015. 119. Stolyarov, I. P.; Demina, L. I.; Cherkashina, N. V. Russ. J. Inorg. Chem. 2011, 56, 1532–1537. 120. Bakhmutov, V. I.; Berry, J. F.; Cotton, F. A.; Murillo, C. A. Dalton Trans. 2005, 1989–1992.

Impurities in Organometallic Catalysis

657

121. Bajwa, S. E.; Storr, T. E.; Hatcher, L. E.; Williams, T. J.; Baumann, C. G.; Whitwood, A. C.; Allan, D. R.; Teat, S. J.; Raithby, P. R.; Fairlamb, I. J. S. Chem. Sci. 2012, 3, 1656–1661. 122. Bedford, R. B.; Bowen, J. G.; Davidson, R. B.; Haddow, M. F.; Seymour-Julen, A. E.; Sparkes, H. A.; Webster, R. L. Angew. Chem. Int. Ed. 2015, 54, 6591–6594. 123. For a review, see: Carole, W. A.; Colacot, T. J. Chem. A Eur. J. 2016, 22, 7686–7695. 124. Carole, W. A.; Bradley, J.; Sarwar, M.; Colacot, T. J. Org. Lett. 2015, 17, 5472–5475. 125. For an overview, see: Grela, K. Ed. Olefin Metathesis: Theory and Practice, Wiley, Weinheim, Germany, 2014. 126. For a review, see: Ogba, O. M.; Warner, N. C.; O’Leary, D. J.; Grubbs R. H. Chem. Soc. Rev. 2018, 47, 4510–4544. 127. For selected examples, see: (a) Wang, H.; Goodman, S. N.; Dai, Q.; Stockdale, G. W.; Clark, W. M. Org. Process Res. Dev. 2008, 12, 226–234.(b) Stark, A.; Ajam, M.; Green, M.; Raubenheimer, H. G.; Ranwell, A.; Ondruschka, B. Adv. Synth. Catal. 2006, 348, 1934–1941; (c) Fu, G.; Grubbs, R. H. J. Am. Chem. Soc. 1992, 114, 7324–7325; (d) Fu, G.; Grubbs, R. H. J. Am. Chem. Soc. 1992, 114, 5426–5427 128. Choi, J.; Kim, H.; Do, T.; Moon, J.; Choe, Y.; Kim, J. G.; Bang, J. J. Polym. Sci. A Polym. Chem. 2019, 57, 726–737. 129. Bielawski, C. W.; Benitez, D.; Grubbs, R. H. J. Am. Chem. Soc. 2003, 125, 8424–8425. 130. For literature precedent in the case of rhenium complexes, see: (a) Kiel, W. A.; Lin, G.-Y.; Gladysz, J. A. J. Am. Chem. Soc. 1980, 102, 3299–3301; (b) Kiel, W. A.; Lin, G.-Y.; Constable, A. G.; McCormick, F. B.; Strause, C. E.; Eisenstein, O.; Gladysz, J. A. J. Am. Chem. Soc. 1982, 104, 4865–4878; (c) Kiel, W. A.; Lin, G.-Y.; Bodner, G. S.; Gladysz, J. A. J. Am. Chem. Soc. 1983, 105, 4958–4972; (d) Kiel, W. A.; Buhro, W. E.; Gladysz, J. A. Organometallics 1994, 3, 879–886 131. For an overview of the field, see: (a) Li, J. J. Ed. C-H Bond Activation in Organic Synthesis, CRC Press, Boca Raton FL, USA, 2017; (b) Crabtree, R. H.; Lei, A. Chem. Rev. 2017, 117, 8481–8482.. and reviews therein; (c) In C-H Bond Activation and Catalytic Functionalization; Dixneuf, P. H., Doucet, H., Eds.; Topics in Organometallic Chemistry Springer: Berlin, Germany, 2016; vol. 56; (d) Roudesly, F.; Oble, J.; Poli, G. J. Mol. Catal. A Chem. 2016, 426, 275–296 132. (a) Gandeepan, P.; Müller, T.; Zell, D.; Cera, G.; Warratz, S.; Ackermann, L. Chem. Rev. 2019, 119, 2192–2452; (b) Singh, K. S. Catalysts 2019, 9, 173; (c) Nareddy, P.; Jordan, F.; Szostak, M. ACS Catal. 2017, 7, 5721–5745; (d) Zha, G.-F.; Qin, H.-L.; Kantchev, E. A. B. RSC Adv. 2016, 6, 30875–30885; (e) Ackermann, L.; Vicente, R. Top. Curr. Chem. 2010, 292, 211–229. 133. Ouellet, S. G.; Roy, A.; Molinaro, C.; Angelaud, R.; Marcoux, J.-F.; O’Shea, P. D.; Davies, I. W. J. Org. Chem. 2011, 76, 1436–1439. 134. (a) Ozdemir, I.; Demir, S.; Çetinkaya, B.; Gourlaouen, C.; Maseras, F.; Bruneau, C.; Dixneuf, P. H. J. Am. Chem. Soc. 2008, 130, 1156–1157; (b) Ackermann, L. Org. Lett. 2005, 7, 3123–3125; (c) Oi, S.; Sakai, K.; Inoue, Y. Org. Lett. 2005, 7, 4009–4011. 135. Ackermann, L. Chem. Rev. 2011, 111, 1315–1345. 136. (a) Arockiam, P.; Poirier, V.; Fischmeister, C.; Bruneau, C.; Dixneuf, P. H. Green Chem. 2009, 11, 1871–1875; (b) Požgan, F.; Dixneuf, P. H. Adv. Synth. Catal. 2009, 351, 1737–1743; (c) Ackermann, L.; Mulzer, M. Org. Lett. 2008, 10, 5043–5045. 137. (a) Drev, M.; Grošelj, U.; Ledinek, B.; Perdih, F.; Svete, J.; Štefane, B.; Požgan, F. ChemCatChem 2018, 10, 3824–3832; (b) Graux, L. V.; Giorgi, M.; Buono, G.; Clavier, H. Dalton Trans. 2016, 45, 6491–6502; (c) Ackermann, L.; Pospech, J.; Potukuchi, H. K. Org. Lett. 2012, 14, 2146–2149; (d) Ackermann, L.; Vicente, R.; Potukuchi, H. K.; Pirovano, V. Org. Lett. 2010, 12, 5032–5035. 138. (a) Frohs, W., Jaeger, H., Eds.; In Industrial Carbon and Graphite Materials: Raw Materials, Production and Applications; Wiley-VCH: Weinheim, Germany, 2021; (b) Kharissova, O. V., Kharisov, B., Eds.; In All-Carbon Composites and Hybrids; Royal Society of Chemistry: Cambridge, UK, 2021. 139. For background on carbon nanotubes, see: (a) Martynková, G. S. Carbon Nanomaterials, De Gruyter STEM, Berlin, Germany, 2021; (b) O’Connell, M. J. Carbon Nanotubes: Properties and Applications; CRC Press: Boca Raton FL, USA, 2018. 140. For background on graphene, see: (a) Foa Torres, L. E. F.; Roche, S.; Charlier, J.-C. Introduction to Graphene-Based Nanomaterials: From Electronic Structure to Quantum Transport, Cambridge University Press, Cambridge, UK, 2020; (b) Johnson, L.; Meany, J. E. Graphene: The Superstrong, Superthin, and Superversatile Material That Will Revolutionize the World; Prometheus: Buffalo, New York, USA, 2018. 141. Takai, K.; Tsujimura, S.; Kang, F.; Inagaki, M. Graphene: Preparations, Properties, Applications, and Prospects; Elsevier: Dordrecht, Netherlands, 2019. 142. Sadjadi, S., Ed.; In Emerging Carbon Materials for Catalysis; Elsevier: Dordrecht, Netherlands, 2020. 143. Dai, L., Ed.; In Carbon-Based Metal-Free Catalysts: Design and Applications; Wiley-VCH: Weinheim, Germany, 2018. 144. For reviews, see: (a) Zhou, Y.; Chen, G.; Zhang, J. J. Mater. Chem. A 2020, 8, 20849–20869; (b) Zhang, L.; Lin, C.-Y.; Zhang, D.; Gong, L.; Zhu, Y.; Zhao, Z.; Xu, Q.; Li, H.; Xia, Z. Adv. Mater. 2019, 31, , 1805252; (c) Liu, D.; Dai, L.; Lin, X.; Chen, J.-F.; Zhang, J.; Feng, X.; Müllen, K.; Zhu, X.; Dai, S. Adv. Mater. 2019, 31, 1804863; (d) Liu, X.; Dai, L. Nat. Rev. Mater. 2016, 1, 16064; (e) Hu, C.; Dai, L. Angew. Chem. Int. Ed. 2016, 55, 11736–11758. 145. (a) Wang, X.-D.; Vinodgopal, K.; Dai, G.-P. Synthesis of carbon nanotubes by catalytic chemical vapor deposition. In Perspective of Carbon Nanotubes; Hosam El-Din Saleh, H. E.-D., El-Sheikh, S. M. M., Eds.; IntechOpen, 2019; (b) Hussein, F. H.; Abdulrazzak, F. H. In Nanomaterials: Biomedical, Environmental, and Engineering Applications; Kanchi, S., Ahmed, S., Sabela, M. I., Hussain, C. M., Eds.; Wiley: Weinheim, Germany, 2018; pp 105–132. Ch. 4. 146. Lin, L.; Deng, B.; Sun, J.; Peng, H.; Liu, Z. Chem. Rev. 2018, 118, 9281–9343. 147. (a) Pumera, M.; Ambrosi, A.; Chng, E. L. K. Chem. Sci. 2012, 3, 3347–3355; (b) Banhart, F. Nanoscale 2009, 1, 201–213; (c) Pumera, M. Langmuir 2007, 23, 6453–6458. 148. Pumera, M., in Carbon Nanotubes: New Research, Ottenhouse, A. P. Ed., Nova Science Publishers, Inc. Hauppauge, NY, USA, 2009; Ch. 2, pp. 75–79. 149. Patnaik, P. Dean’s Analytical Chemistry Handbook, 2nd ed.; McGraw-Hill Professional: NY, USA, 2004. Ch. 1. 150. Kicinski, W.; Dyjak, S. Carbon 2020, 168, 748–845. 151. Tan, S. M.; Pumera, M. ACS Nano 2019, 13, 2681–2728. 152. For a review, see: Pumera, M. ACS Catal. 2020, 10, 7087–7092. 153. Wang, L.; Pumera, M. Appl. Mater. Today 2016, 5, 134–141. 154. Pumera, M.; Iwai, H. Chem. Asian J. 2009, 4, 554–560. 155. Banks, C. E.; Crossley, A.; Salter, C.; Wilkins, S. J.; Compton, R. G. Angew. Chem. Int. Ed. 2006, 45, 2533–2537. 156. (a) Kruusma, J.; Mould, N.; Jurkschat, K.; Crossley, A.; Banks, C. E. Electrochem. Commun. 2007, 9, 2330–2333; (b) Sljukic, B.; Banks, C. E.; Compton, R. G. Nano Lett. 2006, 6, 1556–1558. 157. Ambrosi, A.; Pumera, M. Chem. A Eur. J. 2010, 16, 1786–1792. 158. (a) Batchelor-McAuley, C.; Wildgoose, G. G.; Compton, R. G.; Shao, L.; Green, M. L. H. Sens. Actuators B 2008, 132, 356–360; (b) Dai, X.; Wildgoose, G. G.; Compton, R. G. Analyst 2006, 131, 901–906. 159. (a) Wang, L.; Sofer, Z.; Pumera, M. ACS Nano 2020, 14, 21–25; (b) Mazánek, V.; Luxa, J.; Matejkova, S.; Kucera, J.; Sedmidubsy, S.; Pumera, M.; Sofer, Z. ACS Nano 2019, 13, 1574–1582; (c) Lum, Y.; Kwon, Y.; Lobaccaro, P.; Chen, L.; Clark, E. L.; Bell, A. T.; Ager, J. W. ACS Catal. 2016, 6, 202–209; (d) Chua, C. K.; Sofer, Z.; Khezri, B.; Webster, R. D.; Pumera, M. Phys. Chem. Chem. Phys. 2016, 18, 17875–17880; (e) Lupina, G.; Kitzmann, J.; Costina, I.; Lukosius, M.; Wenger, C.; Wolff, A.; Vaziri, S.; Stling, M. O.; Pasternak, I.; Krajewska, A.; Strupinski, W.; Kataria, S.; Gahoi, A.; Lemme, M. C.; Ruhl, G.; Zoth, G.; Luxenhofer, O.; Mehr, W. ACS Nano 2015, 9, 4776–4785; (f ) Ambrosi, A.; Pumera, M. Nanoscale 2014, 6, 472–476; (g) Chua, C. K.; Ambrosi, A.; Sofer, Z.; Mackova, A.; Havranek, V.; Tomandl, I.; Pumera, M. Chem. A Eur. J. 2014, 20, 15760–15767; (h) Wang, L.; Ambrosi, A.; Pumera, M. Angew. Chem. Int. Ed. 2013, 52, 13818–13821. 160. For background, see: (a) Dean, J. R. Practical Inductively Coupled Plasma Spectrometry, Wiley-VCH, Weinheim, Germany, 2nd ed., 2019.(b) Thomas, R. Practical Guide to ICPMS: A Tutorial for Beginners, 3rd ed.; CRC Press: Boca Raton FL, USA, 2013. 161. Gruber, A. S.; Pozebon, D.; Monteiro, A. L. Tetrahedron Lett. 2001, 42, 7345–7348. 162. See, for example: (a) Petrov, P.; Rajamäki, T.; Corns, W. T.; Goenaga-Infante, H. Atmos. Environ.: X 2020, 8, 100090; (b) He, M.-Y.; Deng, L.; Lu, H.; Jin, Z.-D. J. Anal. At. Spectrom 2019, 34, 1026–1032; (c) Wang, Y.; Brindle, I. D. J. Anal. At. Spectrom 2014, 29, 1904–1911; (d) Li, Y.; Chen, C.; Li, B.; Sun, J.; Wang, J.; Gao, Y.; Zhao, Y.; Chai, Z. J. Anal. At. Spectrom 2006, 21, 94–96; (e) Harrington, C. F.; Merson, S. A.; D’ Silva, T. M. Anal. Chim. Acta 2004, 505, 247–254.

658 163. 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. 191. 192.

193. 194.

195. 196. 197. 198. 199. 200. 201. 202. 203. 204. 205. 206. 207.

208.

Impurities in Organometallic Catalysis For background, see: (a) Grden, M. J. Radioanal. Nucl. Chem. 2020, 323, 677–714; (b) Balaram, V. Trends Anal. Chem. 2016, 80, 83–95 Tu, S.; Yusuf, S.; Muehlfeld, M.; Bauman, R.; Vanchura, B. Org. Process Res. Dev. 2019, 23, 2175–2180. Jones, C. P.; Jurkschat, K.; Crossley, A.; Compton, R. G.; Riehl, B. L.; Banks, C. E. Langmuir 2007, 23, 9501–9504. For a review, see: Vejpravova, J.; Pacakova, B.; Kalbac, M. Analyst 2016, 141, 2639–2656. Kolodiazhnyi, T.; Pumera, M. Small 2008, 4, 1476–1484. For representative examples, see: (a) Bittova, B. P.; Kalbac, M.; Kubickova, S.; Mantlikova, A.; Mangold, S.; Vejpravova, J. Phys. Chem. Chem. Phys. 2013, 15, 5992–6000; (b) Pacakova, B.; Kominkova, Z.; Vejpravova, J.; Mantlikova, A.; Kalbac, M. J. Mater. Sci. 2015, 50, 2544–2553. For background, see: Wang, B.; Anslyn, E. V. Eds. Chemosensors: Principles, Strategies, and Applications, Wiley, New York, USA, 2011. For reviews, see: (a) Upadhyay, S.; Singh, A.; Shivangi, R. S.; Kiran, O.; Tu, N. J. Mol. Struct. 2019, 1193, 89–102.(b) Kwon, N.; Hu, Y.; Yoon, J. ACS Omega 2018, 3, 13731–13751; (c) Wu, D.; Sedgwick, A. C.; Gunnlaugsson, T.; Akkaya, E. U.; Yoon, J.; James, T. D. Chem. Soc. Rev. 2017, 46, 7105–7123. For reviews, see: (a) Balamurugan, R.; Liu, J.-H.; Liu, B.-T. Coord. Chem. Rev. 2018, 376, 196–224; (b) Koide, K. In New Trends in Cross-Coupling: Theory and Applications; Colacot, T., Ed.; Royal Society of Chemistry: Cambridge, UK, 2014. Song, F.; Garner, A. L.; Koide, K. J. Am. Chem. Soc. 2007, 129, 12354–12355. Garner, A. L.; Koide, K. Chem. Commun. 2009, 86–88. Williams, J. M.; Koide, K. Ind. Eng. Chem. Res. 2013, 52, 8612–8615. Song, F.; Carder, E. J.; Kohler, C. C.; Koide, K. Chem. A Eur. J. 2010, 16, 13500–13508. Bu, X.; Koide, K.; Carder, E. J.; Welch, C. J. Org. Process Res. Dev. 2013, 17, 108–113. Koide, K.; Tracey, M. P.; Bu, X.; Jo, J.; Williams, M. J.; Welch, C. J. Nat. Commun. 2016, 7, 2–8. Lukomski, L.; Pohorilets, I.; Koide, K. Org. Process Res. Dev. 2020, 24, 85–95. Kumari, N.; Dey, N.; Kumar, K.; Bhattacharya, S. Chem. Asian J. 2014, 9, 3174–3181. Li, D.; Campbell, L. D.; Austin, B. A.; Koide, K. ChemPlusChem 2012, 281–283. Inamoto, K.; Campbell, L. D.; Doi, T.; Koide, K. Tetrahedron Lett. 2012, 53, 3147–3148. Inamoto, K.; Hasegawa, C.; Hiroya, K.; Kondo, Y.; Osako, T.; Uozumi, Y.; Doi, T. Chem. Commun. 2012, 48, 2912–2914. For reviews covering claims of “metal-free” reactions, see: (a) Piazzolla, F.; Colognese, F.; Temperini, A. Curr. Org. Chem. 2018, 22, 2537–2554; (b) Qin, Y.; Zhu, L.; Luo, S. Chem. Rev. 2017, 117, 9433–9520; (c) Rick, R. A.; Lin, C. S. K.; Vargas, C.; Luque, R. Org. Biomol. Chem. 2014, 12, 10–35; (d) Mehta, V. P.; Punji, B. RSC Adv. 2013, 11957–11986 Pentsak, E. O.; Eremin, D. B.; Gordeev, E. G.; Ananikov, V. P. ACS Catal. 2019, 9, 3070–3081. Baars, H.; Beyer, A.; Kohlhepp, S. V.; Bolm, C. Org. Lett. 2014, 16 (2), 536–539. Pichette Drapeau, M.; Fabre, I.; Grimaud, L.; Ciofini, I.; Ollevier, T.; Taillefer, M. Angew. Chem. Int. Ed. 2015, 54, 10587–10591. Reeves, B. M.; Hepburn, H. B.; Grozavu, A.; Lindsay-Scott, P. J.; Donohoe, T. J. Angew. Chem. Int. Ed. 2019, 58, 15697–15701. Marien, N.; Reddy, B. N.; De Vleeschouwer, F.; Goderis, S.; Van Hecke, K.; Verniest, G. Angew. Chem. Int. Ed. 2018, 57, 5660–5664. Casotti, G.; Ciancaleoni, G.; Lipparini, F.; Nieri, C.; Iuliano, A. Chem. Sci. 2020, 11, 257–263. Yanagisawa, S.; Itami, K. ChemCatChem 2011, 3, 827–829. Yanagisawa, S.; Ueda, K.; Taniguchi, T.; Itami, K. Org. Lett. 2008, 10, 4673–4676. For context of using scavengers to remove palladium from reaction mixtures, see: (a) Yamada, T.; Matsuo, T.; Ogawa, A.; Ichikawa, T.; Kobayashi, Y.; Masuda, H.; Miyamoto, R.; Bai, H. Z.; Meguro, K.; Sawama, Y.; Monguchi, Y.; Sajiki, H. Org. Process Res. Dev. 2019, 23, 462–469; (b) Yamada, T.; Kobayashi, Y.; Ito, N.; Ichikawa, T.; Park, K.; Kunishima, K.; Ueda, S.; Mizuno, M.; Adachi, T.; Sawama, Y.; Monguchi, Y.; Sajiki, H. ACS Omega 2019, 4, 10243–10251; (c) Ren, H.; Strulson, C. A.; Humphrey, G.; Xiang, R.; Li, G. T.; Gauthier, D. R.; Maloney, K. M. Green Chem. 2017, 19, 4002–4006; (d) Miyamoto, H.; Sakumoto, C.; Takekoshi, E.; Maeda, Y.; Hiramoto, N.; Itoh, T.; Kato, Y. Org. Process Res. Dev. 2015, 19, 1054–1061; (e) Mondal, B.; Wilkes, R. D.; Percy, J. M.; Tuttle, T.; Black, R. J. G.; North, C. Dalton Trans. 2014, 43, 469–478; (f ) Reginato, G.; Sadler, P.; Wilkes, R. D. Org. Process Res. Dev. 2011, 15, 1396–1405; (g) Wang, L.; Green, L.; Li, Z.; McCabe Dunn, J.; Bu, X.; Welch, C. J.; Li, C.; Wang, T.; Tu, Q.; Bekos, E.; Richardson, D.; Eckert, J.; Cui, J. Org. Process Res. Dev. 2011, 15, 1371–1376; (h) Magano, J. In Transition Metal-Catalyzed Couplings in Process Chemistry: Case Studies from the Pharmaceutical Industry; Magano, J., Dunetz, J. R., Eds.; Wiley-VCH: Weinheim, Germany, 2013; pp 313–355; (i) Mendonca, T. In The Power of Functional Resins in Organic Synthesis; Tulla-Puche, J., Albericio, F., Eds.; Wiley-VCH: Weinheim, Germany, 2008; pp 227–243. Vallée, F.; Mousseau, J. J.; Charette, A. B. J. Am. Chem. Soc. 2010, 132, 1514–1516. For reviews, see: (a) Zweig, J. E.; Kim, D. E.; Newhouse, T. R. Chem. Rev. 2017, 117, 11680–11752.(b) Sivanandan, S. T.; Shaji, A.; Ibnusaud, I.; Seechurn, C. C. C. J.; Colacot, T. J. Eur. J. Org. Chem. 2015, 38–49; (c) Johansson, C. C. C.; Colacot, T. J. Angew. Chem. Int. Ed. 2010, 49, 676–707; (d) Bellina, F.; Rossi, R. Chem. Rev. 2010, 110, 1082–1114 Sun, C.-L.; Li, H.; Yu, D.-G.; Yu, M.; Zhou, X.; Lu, X.-Y.; Huang, K.; Zheng, S.-F.; Li, B.-J.; Shi, Z.-J. Nat. Chem. 2010, 2, 1044–1049. Yamamoto, E.; Maeda, S.; Taketsugu, T.; Ito, H. Synlett 2017, 28, 1258–1267. Yamamoto, E.; Izumi, K.; Horita, Y.; Ito, H. J. Am. Chem. Soc. 2012, 134, 19997–20000. Crabtree, R. H. Chem. Rev. 2012, 112, 1536–1554. Rosenberg, B. Platin. Met. Rev. 1973, 15, 42–51. (a) Dralle Mjos, K.; Orvig, C. Chem. Rev. 2014, 114, 4540–4563; (b) Alderden, R. A.; Hall, M. D.; Hambley, T. W. J. Chem. Educ. 2006, 83, 728–734. (a) Ghosh, S. Bioorg. Chem. 2019, 88, 102925; (b) Aldossary, S. A. Biomed. Pharmacol. J. 2019, 12, 7–15. (a) Dabrowiak, J. C. Metals in Medicine; John Wiley & Sons: Hoboken, New Jersey, USA, 2017; (b) In Cisplatin: Chemistry and Biochemistry of a Leading Anticancer Drug; Lippert, B., Ed.; John Wiley & Sons: Hoboken, New Jersey, USA, 2007. Reddy, M. M.; Reddy, K. H.; Reddy, M. U. Pharmaceut. Reg. Affairs 2016, 5, 168. See, for example: (a) Ho, W. C.; Chung, K.; Ingram, A. J.; Waymouth, R. M. J. Am. Chem. Soc. 2018, 140, 748–757; (b) Albers, P.; Pietsch, J.; Parker, S. F. J. Mol. Catal. A Chem. 2001, 173, 275–286 Xu, L.; Liu, F.-Y.; Zhang, Q.; Chang, W.-J.; Liu, Z. L.; Lv, Y.; Yu, H.-Z.; Xu, J.; Dai, J.-J.; Xu, H. J. Nat. Catal. 2021, 4, 71–78. Novák, Z.; Adamik, R.; Csenki, J. T.; Béke, F.; Gavaldik, R.; Varga, B.; Nagy, B.; May, Z.; Daru, J.; Gonda, Z.; Tolnai, G. L. ChemRxiv 2021. https://doi.org/10.26434/ chemrxiv.14071247.v1. Avanthay, M.; Bedford, R.; Begg, C.; Böse, D.; Clayden, J.; Davis, S.; Eloi, J.-C.; Goryunov, G. P.; Hartung, I. V.; Heeley, J.; Khaikin, K. A.; Kitching, M.; Krieger, J.; Kulyabin, P. S.; Lennox, A.; Nolla-Saltiel, R.; Pridmore, N. E.; Rowsell, B. J. S.; Sparkes, H. A.; Uborsky, D. V.; Voskoboynikov, A. Z.; Walsh, M.; Wilkinson, H. J. ChemRxiv 2021. https://doi.org/10.26434/chemrxiv.14237288.v1. Sun, C. L.; Shi, Z. J. Chem. Rev. 2014, 114, 9219–9280.