Synthetic Diamond Films: Preparation, Electrochemistry, Characterization, and Applications [1 ed.] 0470487585, 9780470487587

Citation preview

Synthetic Diamond Films

Synthetic Diamond Films Preparation, Electrochemistry, Characterization, and Applications

Edited by

Enric Brillas Carlos Alberto Mart´ınez-Huitle

The Wiley Series on Electrocatalysis and Electrochemistry Series Editor:

Andrzej Wieckowski

A John Wiley & Sons, Inc., Publication

Copyright  2011 by John Wiley & Sons, Inc. All rights reserved Published by John Wiley & Sons, Inc., Hoboken, New Jersey Published simultaneously in Canada No part of this publication may be reproduced, stored in a retrieval system, or transmitted in any form or by any means, electronic, mechanical, photocopying, recording, scanning, or otherwise, except as permitted under Section 107 or 108 of the 1976 United States Copyright Act, without either the prior written permission of the Publisher, or authorization through payment of the appropriate per-copy fee to the Copyright Clearance Center, Inc., 222 Rosewood Drive, Danvers, MA 01923, (978) 750-8400, fax (978) 750-4470, or on the web at www.copyright.com. Requests to the Publisher for permission should be addressed to the Permissions Department, John Wiley & Sons, Inc., 111 River Street, Hoboken, NJ 07030, (201) 748-6011, fax (201) 748-6008, or online at http://www.wiley.com/go/permission. Limit of Liability/Disclaimer of Warranty: While the publisher and author have used their best efforts in preparing this book, they make no representations or warranties with respect to the accuracy or completeness of the contents of this book and specifically disclaim any implied warranties of merchantability or fitness for a particular purpose. No warranty may be created or extended by sales representatives or written sales materials. The advice and strategies contained herein may not be suitable for your situation. You should consult with a professional where appropriate. Neither the publisher nor author shall be liable for any loss of profit or any other commercial damages, including but not limited to special, incidental, consequential, or other damages. For general information on our other products and services or for technical support, please contact our Customer Care Department within the United States at (800) 762-2974, outside the United States at (317) 572-3993 or fax (317) 572-4002. Wiley also publishes its books in a variety of electronic formats. Some content that appears in print may not be available in electronic formats. For more information about Wiley products, visit our web site at www.wiley.com. Library of Congress Cataloging-in-Publication Data: Synthetic diamond films : preparation, electrochemistry, characterization, and applications / edited by Enric Brillas, Carlos Alberto Mart´ınez-Huitle. p. cm.—(The Wiley series on electrocatalysis and electrochemistry) Includes index. ISBN 978-0-470-48758-7 (cloth) 1. Diamonds—Electric properties. 2. Diamond thin films. I. Brillas, Enric. II. Mart´ınez-Huitle, Carlos Alberto. TK7871.15.D53S96 2011 666′ .88—dc22 2010053392 Printed in Singapore oBook ISBN: 978-1-118-06236-4 ePDF ISBN: 978-1-118-06234-0 ePub ISBN: 978-1-118-06235-7 10 9 8 7 6 5 4 3 2 1

Contents

PREFACE

xix

PREFACE TO THE WILEY SERIES ON ELECTROCATALYSIS AND ELECTROCHEMISTRY

xxiii

CONTRIBUTORS

xxv

PART I

SYNTHESIS OF DIAMOND FILMS

1. Electrochemistry on Diamond: History and Current Status

1 3

John C. Angus

1.1

Enabling Technologies / 3 1.1.1 Chemical Vapor Deposition of Diamond / 3 1.1.2 Doping of Diamond / 4 1.1.3 Surface Characterization of Diamond / 5

1.2

First Studies of the Electrochemistry on Diamond / 5 1.2.1 From 1987 to 1996 / 5 1.2.2 From 1996 to Present / 6

1.3

Development of Electrochemical Applications of Diamond / 8 1.3.1 Surface Functionalization / 8 1.3.2 Destruction of Wastes / 9 1.3.3 Sensors and Electroanalysis / 9

1.4

Other Directions / 10 1.4.1 Biolectronic Applications / 10 1.4.2 Anomalous Surface Conductivity of Diamond / 11 v

vi

CONTENTS

1.5

Conclusions / 13 References / 13

2. Synthesis of Diamond Films

21

Vadali V. S. S. Srikanth and Xin Jiang

2.1

Introduction / 21

2.2

Diamond Film CVD Techniques / 23 2.2.1 History / 23 2.2.2 Thermal Decomposition Techniques / 25 2.2.2.1 Hot Filament Chemical Vapor Deposition (HFCVD) / 25 2.2.2.2 Oxy-Acetylene Torch Method / 25 2.2.3 Plasma-Aided Deposition Techniques / 26 2.2.3.1 Microwave Plasma-Enhanced CVD (MWCVD) / 26 2.2.3.2 DC Plasma CVD / 27 2.2.3.3 RF Plasma CVD / 28 2.2.3.4 Electron Cyclotron Resonance Microwave Plasma-Assisted CVD / 28

2.3

Diamond Nucleation and Growth / 28 2.3.1 Nucleation / 28 2.3.1.1 Definition and Types / 28 2.3.1.2 Methods / 30 2.3.2 Growth / 32 2.3.3 Role of Hydrogen and Oxygen / 39

2.4

Diamond Epitaxy / 40

2.5

Nanodiamond Thin Films / 45

2.6

Diamond Nanocomposite Films / 46

2.7

Conclusions / 48 References / 48

3. Types of Conducting Diamond Materials and Their Properties Marco A. Quiroz and Erick R. Bandala

3.1

Introduction / 57

3.2

Conducting Diamond Materials (CDMs) / 62

3.3

CDM Preparation Procedures / 63

3.4

CDM Doping Materials / 63 3.4.1 Characteristics of Boron-Doped CDMs / 63 3.4.2 Electrochemical Properties / 64

57

vii

CONTENTS

3.4.3 3.4.4 3.4.5 3.4.6

Photoelectrochemical Properties / 66 Optical Spectroscopy Properties / 67 Photo- and Cathodoluminescence Properties / 68 Electrical Conductivity and Superconductivity Properties / 69

3.5

Non-Boron-Doped CDMs / 69

3.6

Conclusions / 71 References / 71

PART II ELECTROCHEMISTRY OF DIAMOND FILMS

77

4. Electrochemistry of Diamond

79

Yuri Pleskov

4.1

Introduction / 79

4.2

Principal Electrochemical Properties of Diamond / 80

4.3

The Effect of Semiconductor Nature of Diamond on its Electrochemical Behavior / 83

4.4

The Effect of Crystal Structure on the Electrochemical Behavior of Diamond / 92 4.4.1 The Effect of Crystallographic Orientation of Crystal Faces / 92 4.4.2 The Effect of Surface Morphology / 95 4.4.3 The Effect of the Diamond Grain Size (or the Film Thickness, or the sp 2 -Carbon Impurity) / 98

4.5

Diamond-Based Nanostructures as Electrode Materials: Vacuum-Annealed Undoped Polycrystalline Diamond / 102

4.6

Conclusions / 106

4.7

Acknowledgments / 106 References / 106

5. Applications of Polycrystalline and Modified Functional Diamond Electrodes Yasuaki Einaga and Akira Fujishima

5.1

Introduction / 109

5.2

Preparation of BDD Electrodes / 110

5.3

Electrochemical Properties of BDD as Electrode Materials / 111

5.4

Applications in Electrochemical Analysis Using Polycrystalline BDD electrodes / 111 5.4.1 Detection of Free Chlorine / 111

109

viii

CONTENTS

5.4.2 5.4.3

Detection of Oxalic Acid / 113 Proteins (Including Cancer Markers) / 113

5.5

Modified Functional BDD Electrodes / 116 5.5.1 Production of High-Concentration Ozone-Water Using Free-Standing Perforated Diamond / 116 5.5.2 Modified Functional BDD Electrodes for Electrochemical Analysis / 119 5.5.2.1 Ion-Implanted BDD Electrodes / 119 5.5.2.2 Selective Detection of As(III) and As(V) by Stripping Voltammetry / 124 5.5.2.3 In vivo Dopamine Detection by BDD Microelectrodes / 125 5.5.2.4 BDD Nanograss Array (Whisker BDD) / 126

5.6

Conclusions / 130

5.7

Acknowledgments / 130 References / 131

6. Diamond Ultramicroelectrodes and Nanostructured Electrodes

133

Katherine B. Holt

6.1

Introduction / 133

6.2

Ultramicroelectrodes: Definition and Electrochemical Characteristics / 134

6.3

Boron-Doped Diamond UMEs / 136 6.3.1 Substrate Preparation and Growth of Diamond Films / 136 6.3.2 Insulation Methods and Control of Exposed Electrode Geometry / 140 6.3.3 Electrochemical Performance and Applications / 142

6.4

Boron-Doped Diamond UME Arrays / 143 6.4.1 Fabrication of BDD UME Arrays / 144 6.4.2 Electrochemical Performance and Applications / 146

6.5

Nanostructured BDD Electrodes / 147 6.5.1 Random Array BDD Nanodisk Electrodes / 147 6.5.2 Fabrication of Nanostructured BDD Arrays / 148 6.5.3 Electrochemical Performance and Applications of Nanostructured BDD Electrodes / 149

6.6

Conclusions and Future Directions / 150 References / 151

CONTENTS

ix

PART III ELECTROANALYTICAL APPLICATIONS

153

7. Electroanalytical Applications of Diamond Films

155

Weena Siangproh, Amara Apilux, Pimkwan Chantarateepra, and Orawon Chailapakul

7.1

Introduction / 155

7.2

Pharmaceutical Compounds / 156

7.3

Biomolecules or Biological Compounds / 159

7.4

Pollutant Compounds / 162

7.5

Heavy Metals / 165

7.6

Food and Dietary Contaminants / 166

7.7

Miscellaneous / 168

7.8

Conclusions / 170

7.9

Acknowledgments / 178 References / 178

8. Cathodic Pretreatment of Boron-Doped Diamond Electrodes and Their Use in Electroanalysis Leonardo S. Andrade, Giancarlo R. Salazar-Banda, Romeu C. Rocha-Filho, and Orlando Fatibello-Filho

8.1

Introduction / 181

8.2

Cathodic Pretreatment of Conductive Diamond Films / 182

8.3

Electroanalytical Applications / 192 8.3.1 General Aspects / 192 8.3.2 Determination of Pesticides in Environmental Samples / 193 8.3.2.1 Carbaryl / 193 8.3.2.2 4-Nitrophenol / 193 8.3.2.3 Chlorophenols / 196 8.3.3 Determination of Substances in Food Samples / 198 8.3.3.1 Aspartame / 198 8.3.3.2 Sodium Cyclamate / 199 8.3.3.3 Aspartame and Sodium Cyclamate / 200 8.3.3.4 Total Phenols / 201 8.3.4 Determination of Substances in Pharmaceutical Samples / 201 8.3.4.1 Sulfamethoxazole and Trimethoprim / 201 8.3.4.2 Sulfamethoxazole and Sulfadiazine / 205

181

x

CONTENTS

8.3.4.3 8.3.4.4 8.3.4.5 8.3.4.6

Acetylsalicylic Acid / 205 Paracetamol and Caffeine / 206 Sildenafil Citrate (Viagra) / 206 Lidocaine / 207

8.4

Gold Deposition and Stripping / 209

8.5

Conclusions / 209 References / 210

PART IV INDUSTRIAL APPLICATIONS 9. Use of Boron-Doped Diamond Electrode in Electrochemical Generation and Applications of Ferrate

213

215

Virender K. Sharma, Enric Brillas, Ignasi Sir´es, and Karel Bouzek

9.1

Introduction / 215

9.2

Electrochemical Generation of the Ferrate Ion with Iron Anodes / 217

9.3

Electrochemical Generation of the Ferrate Ion with Inert Anodes / 222

9.4

Electrochemical Generation of the Ferrate Ion with Boron-Doped Diamond Anode / 223 9.4.1 Acidic Medium / 223 9.4.2 Alkaline Medium / 225

9.5

Applications / 228 9.5.1 Common Inert Anodes / 228 9.5.2 Iron Anodes / 229 9.5.3 BDD Anode / 230

9.6

Conclusions / 233

9.7

Acknowlegments / 233 References / 233

10. Electrochemical Oxidation of Organic Compounds Induced by Electro-Generated Free Hydroxyl Radicals on BDD Electrodes Agnieszka Kapałka, Helmut Baltruschat, and Christos Comninellis

10.1

Introduction / 237

10.2

Influence of Anode Material on the Reactivity of Electrolytic Hydroxyl Radicals / 238

10.3

Electro-Generation and Detection of Quasi-Free Hydroxyl Radicals on BDD Electrode / 240

237

xi

CONTENTS

10.3.1 10.3.2 10.3.3 10.3.4

Hydroxyl Radicals Spin Trapping / 240 Trapping by Salicylic Acid / 240 Competitive Reactions / 242 Formation of Hydrogen Peroxide / 242

10.4

Concentration Profile of Hydroxyl Radicals on BDD Electrode / 244 10.4.1 HO • Concentration Profile during Oxygen Evolution / 244 10.4.2 HO • Concentration Profile during Electro-Oxidation of Organic Compound / 246

10.5

Kinetic Model of Organics Oxidation on BDD Anode / 248 10.5.1 Electrolysis under Current Limited Control (japplied < jlim ) / 249 10.5.2 Electrolysis under Mass Transport Control (japplied > jlim ) / 251

10.6

Electrochemically Induced Mineralization of Organic Compounds by Molecular Oxygen / 253

10.7

Conclusions / 256

10.8

Exercises / 256 10.8.1 Solutions / 257 References / 260

11. Modeling of Electrochemical Process for Water Treatment Using Diamond Films

261

Onofrio Scialdone and Alessandro Galia

11.1

Introduction / 261

11.2

Theoretical Models / 263 11.2.1 General Considerations / 263 11.2.2 Oxidation of Organic Pollutants in Water at BDD by Means of Direct Anodic Oxidation or Reaction with Electro-Generated Hydroxyl Radicals (“Direct Processes”) / 265 11.2.2.1 The Model of Comninellis and Coauthors / 268 11.2.2.2 The Theoretical Works of Polcaro and Coauthors / 272 11.2.2.3 The Approach Proposed by Rodrigo and Coauthors / 273 11.2.3 Oxidation of Organic Pollutants in Water by Means of Electro-Generated Oxidants (“Indirect Processes”) Such as Active Chlorine / 274

11.3

Conclusions / 278

11.4

Acknowledgments / 279 References / 279

xii

CONTENTS

12. Production of Strong Oxidizing Substances with BDD Anodes

281

Ana S´anchez-Carretero, Cristina S´aez, Pablo Ca˜nizares, and Manuel A. Rodrigo

12.1

Electrolyses with Conductive-Diamond Anodes / 281

12.2

Production and Storage of Oxidizing Substances: Experimental Setups / 283

12.3

Production of Hydroxyl Radicals with Conductive-Diamond Anodes / 284

12.4

Synthesis of Peroxoacids and Peroxosalts / 288 12.4.1 Peroxosulphuric Acids / 288 12.4.2 Peroxodiphosphate Salts / 292 12.4.3 Monoperoxophosphoric Acid / 296

12.5

Synthesis of Halogen Oxoanions / 300 12.5.1 Perchlorates / 300 12.5.2 Perbromates / 300

12.6

Synthesis of Ferrates / 301

12.7

Effect of the Type of Diamond on the Efficiency of the Production of Oxidants / 305

12.8

Conclusions / 307

12.9

Acknowledgments / 308 References / 308

13. Ozone Generation Using Boron-Doped Diamond Electrodes

311

Yunny Meas, Luis A. Godinez, and Erika Bustos

13.1

Introduction / 311

13.2

Ozone 13.2.1 13.2.2 13.2.3 13.2.4

13.3

Technologies for Producing Ozone / 317 13.3.1 Corona Discharge Technique / 317 13.3.2 Electrical Discharge Ozone Generators (EDOGs) / 319 13.3.3 Electrolytic Ozone Generators (ELOGs) / 319 13.3.3.1 Anodes for Electrochemically Producing Ozone / 320 13.3.3.2 Boron-Doped Diamond (BDD) / 323

13.4

Reaction Mechanism for the Production of Ozone with Boron-Doped Diamond / 325

/ 311 Physical and Chemical Properties of Ozone / 312 Production of Ozone / 313 Importance of Ozone Applications / 313 Efficiency and Production / 315

CONTENTS

13.5

xiii

Conclusions / 326 References / 327

14. Application of Synthetic Diamond Films to Electro-Oxidation Processes

333

Marco Panizza

14.1

Introduction / 333

14.2

Application in Wastewater Treatment / 335 14.2.1 Oxidation in the Potential Region before Oxygen Evolution / 335 14.2.2 Oxidation in the Potential Region of Oxygen Evolution / 339 14.2.3 Influence of the Nature of Organic Pollutants / 342 14.2.4 Influence of the Concentrations of Organic Compounds / 343 14.2.5 Influence of the Applied Current Density / 343 14.2.6 Influence of the Flow Rate / 345 14.2.7 Influence of the Temperature / 345 14.2.8 Comparison with Other Electrode Materials / 346

14.3

Application in Organic Electrosynthesis / 347

14.4

Conclusions / 348 References / 349

15. Fabrication and Application of Ti/BDD for Wastewater Treatment

353

Xueming Chen and Guohua Chen

15.1

Fabrication of Stable Ti/BBD Electrodes / 353 15.1.1 Introduction / 353 15.1.2 HFCVD Facility / 354 15.1.3 HFCVD Parameter Optimization / 354 15.1.4 Reactive Gas Component Improvement / 357 15.1.5 Methods to Enhance the Service Life of Ti/BDD / 363

15.2

Use of 15.2.1 15.2.2 15.2.3

15.3

Conclusions / 369

Ti/BDD Electrodes for Wastewater Treatment / 365 Oxidation of Acetic and Maleic Acids / 365 Oxidation of Phenol / 365 Oxidation of Dyes / 366

References / 369 16. Application of Diamond Films to Water Disinfection Jessica H. Bezerra Rocha and Carlos A. Mart´ınez-Huitle

16.1

Introduction / 373

373

xiv

CONTENTS

16.2

Disinfection Water / 374

16.3

Science and Technology for Water Purification / 375

16.4

Electrochemical Disinfection/Purification Systems / 376

16.5

Diamond Films for Drinking Water Disinfection / 384

16.6

Production of Inorganic Disinfection by-Products and Inorganic Species Elimination / 388 16.6.1 Chloride, Chlorite, and Chlorate Ions / 389 16.6.2 Perchlorate in Drinking Water / 392 16.6.3 Electrolysis of Nitrates / 394

16.7

Electrochemical Free-Chlorine Systems Using Diamond Films / 396

16.8

Conclusions / 400 References / 400

17. Fenton-Electrochemical Treatment of Wastewaters for the Oxidation of Organic Pollutants Using BDD

405

Enric Brillas

17.1

Introduction / 405

17.2

Fundamentals of Fenton’s Electrochemistry / 406

17.3

Electrogeneration of H2 O2 and Regeneration of Fe2+ / 409

17.4

Degradation of Organics in BDD/O2 Tank Reactors / 413 17.4.1 Herbicides / 414 17.4.2 Dyes / 417 17.4.3 Pharmaceuticals and Amino Acids Precursors / 420

17.5

Degradation of Organics in others Tank Reactors with a BBD Anode / 426

17.6

Degradation of Organics in Batch Recirculation BDD/O2 Flow Cells / 427

17.7

Conclusions / 433 References / 433

18. Electrochemical Energy Storage and Energy Conversion Systems with Diamond Films Juan M. Peralta-Hern´andez, Aracely Hern´andez-Ram´ırez, Jorge L. Guzm´an-Mar, Laura Hinojosa-Reyes, Giancarlo R. Salazar-Banda, and Carlos A. Mart´ınez-Huitle

18.1

Introduction / 437

18.2

Different Techniques Used to Modify BDD Films / 438 18.2.1 Microemulsion Synthesis / 438

437

xv

CONTENTS

18.2.2 18.2.3 18.2.4

Thermal Deposition / 443 Electrodeposition / 446 18.2.3.1 Electrodeposition of Metal Particles on BDD / 449 Sol-Gel Modification / 453

18.3

Application of Modified BDD Films as Electrocatalytic Surfaces for Fuel Cells / 459

18.4

Application of BDD Films in Batteries / 466

18.5

Application of BDD Electrodes as Electrochemical Capacitors / 474

18.6

Conclusions / 477 References / 478

19. Use of Diamond Films in Organic Electrosynthesis

483

Siegfried R. Waldvogel, Axel Kirste, and Stamo Mentizi

19.1

Introduction / 483

19.2

Specific Features of BDD Electrodes / 485

19.3

Stability of BDD Electrodes in Organic Media / 487

19.4

Electrolysis Cells for BDD Electrodes for Organic Transformations / 489

19.5

Anodic 19.5.1 19.5.2 19.5.3 19.5.4 19.5.5 19.5.6

19.6

Cathodic Synthesis on BDD Electrodes / 504 19.6.1 Reduction of Oximes / 504 19.6.2 Reductive Carboxylation / 505

19.7

Conclusions / 506

19.8

Acknowledgment / 506

Transformations on BBD Electrodes / 491 Alkoxylation Reactions / 491 Fluorination Reactions / 493 Cyanation Reactions / 494 Cleavage of C,C-Bonds / 495 Oxidation of Activated Carbon Atoms / 496 Anodic Phenol Coupling Reaction / 496 19.5.6.1 Anodic Homo-Coupling of Phenolic Substrates / 497 19.5.6.2 Nonsymmetrical Phenol Coupling and Phenol-Arene Cross-Coupling Reaction / 499

References / 507

xvi

CONTENTS

PART V BIOELECTROCHEMICAL APPLICATIONS

511

20. Diamond Sensors for Neurochemistry

513

Bhavik Anil Patel

20.1

Introduction / 513

20.2

Central and Peripheral Nervous System / 513

20.3

The Process of Neurotransmission / 514 20.3.1 Neurotransmitters / 516

20.4

Electroanalytical Methods to Study Neurotransmitter Release / 517 20.4.1 Sensors Utilized / 519

20.5

Limitations of Current Techniques for In Vitro and In Vivo Monitoring / 520 20.5.1 Long-Term Recordings / 521 20.5.2 Fouling from Large Biomolecules / 523 20.5.3 Fouling from Redox Reaction By-Products / 525

20.6

Applications of Diamond Sensors and Devices in Neurochemistry / 529 20.6.1 Recording Neuronal Activity / 529 20.6.2 Single Cell Measurements of Vesicular Release / 530 20.6.3 Neurotransmitter Release from Sympathetic Nerves Innervating Mesenteric Arteries / 531 20.6.4 Measuring Transmitter Release from the Gastrointestinal Tract / 533 20.6.4.1 Detection of Histamine Release from Enterochromaffin-Like Cells Located in the Stomach / 534 20.6.4.2 Monitoring Serotonin Release from Enterochromaffin Cells Located in the Mucosa / 535 20.6.4.3 Monitoring Nitric Oxide Release from Myenteric Plexus Neurons / 537 20.6.5 Studying the Neurotransmitter Clearance Process / 538 20.6.5.1 Measurements of Multiple Transmitters from Brain Synaptosomes / 539 20.6.5.2 Investigation of Serotonine Clearance by Transporters Present on Lymphocytes / 540 20.6.6 In vitro and In vivo Measurements from the Central Nervous System / 541 20.6.6.1 In vitro Measurements / 541 20.6.6.2 In vivo Measurements from Anesthetized Animals / 542

20.7

Conclusions and Outlook for the Future / 543

CONTENTS

20.8

xvii

Acknowledgments / 544 References / 544

21. DNA-Modified Diamond Films

551

Nianjun Yang and Christoph E. Nebel

21.1

Introduction / 551

21.2

Diamond Transducer Properties / 558 21.2.1 CVD Diamond Growth / 558 21.2.2 Surface Terminations / 562 21.2.3 Diamond Nanotexture and Wire Formation / 564

21.3

Surface Modification of Diamond / 571 21.3.1 Photochemical Surface Modification of Intrinsic Diamond / 571 21.3.2 Electrochemical Surface Functionalization of Boron-Doped Diamond / 582 21.3.3 Tip Functionalization of Diamond Nanotextures / 589

21.4

DNA Molecules on Diamond / 593 21.4.1 DNA Attachment / 593 21.4.2 Characterizations of DNA Layers / 595

21.5

Sensing 21.5.1 21.5.2 21.5.3

21.6

Summary and Outlook / 612

21.7

Acknowledgments / 614

of DNA Hybridization / 602 DNA Field Effect Transistor / 602 Cyclic Voltammetry and Impedance Spectroscopy / 606 DNA Sensing on Nanotextured Diamond Surfaces / 609

References / 614 INDEX

621

Preface

Diamond is an extremely hard crystalline form of carbon and it is considered an excellent material for many applications due to its unusual physical and chemical properties. For this reason, it has long attracted the attention of scientists and the public. Interest in diamond has been further increased by the discovery of the possibility to produce polycrystalline diamond films with mechanical and electronic properties comparable with natural diamond. Over the last few years, the number of publications has increased considerably regarding the synthesis and/or applications of this new material. Currently, synthetic diamond films have been the subject of applications and fundamental research in several fields of the science. Much effort was spent during the 1960s and 1970s to investigate diamond synthesis until it was successfully achieved by using the chemical vapor deposition (CVD) technique with excellent diamond growth rates, which led to good prospects for films being used for some industrial applications. Since their introduction into electrochemical research in 1987, doped-diamond electrodes have become more and more popular. This is based on their unique properties that distinguish them from conventional electrode materials and make many electrochemical processes more attractive or even possible. These electrodes, then, have been the subject of a large variety of applications and fundamental research in electrochemistry, opening up a novel branch known as the “electrochemistry of synthetic diamond films.” Almost every aspect of electrochemistry has been impacted by the diamond electrode, from instrumental analysis to industrial applications. Electrically conductive films of boron-doped diamond (BDD) have gained increasing popularity in many electrochemical applications, in large part due to the fact that very high quality films possess background currents that can be some orders of magnitude lower than those other types of electrode materials. Other important properties of these electrodes are related to their large potential window, low adsorption, corrosion stability in very aggressive media, high efficiency in oxidation processes, and very low doublelayer capacitance. Therefore, diamond films have become suitable materials for several purposes classified in different areas: synthesis of chemicals, modification of diamond surfaces, electroanalysis, water disinfection, destruction of pollutants in waters, and so xix

xx

PREFACE

on. The versatility of these materials has also been extended to develop sensors, microelectrodes, nanoelectrodes, and biosensors. Recently, synthetic diamond is starting to be commercialized for practical purposes—for example, diamond electrochemical detectors for liquid chromatography and large-scale diamond films for industrial wastewater treatment. The future for the synthetic diamond films is bright. These materials are starting to make important contributions for the measurements and understanding of an extensive range of chemical processes. Several other complementary techniques are emerging and can provide new knowledge for the chemistry of materials. Future interdisciplinary developments based on the close collaboration of chemists, electrochemists, engineers, biologists, and neuroscientists can be envisaged to ensure an effective application and productive use of synthetic diamonds to answer important chemical and medical questions and to resolve vital environmental problems. Several areas of interest have been opened up by these developments, including the measurement of release from single neurons, new sensors that are smaller and yet faster responding, sensor arrays for simultaneously detecting several analytes, novel biosensors for new neurochemical species of interest, and medical and clinical applications to neurochemistry and neuroscience. Synthetic Diamond Films: Preparation, Electrochemistry, Characterization, and Applications is a timely book that has gathered together the best international experts of the electrochemical community and who have imagined a large number of approaches to investigate the electrochemical properties of diamond films and their characteristics. All of these contributing experts now focus on aspects that promote efficiency, selectivity, and high performances for a broad variety of laboratory and industrial diamond applications with goals ranging from organic synthesis to environmental problems. The first part of the book (Chapters 1 to 3) deals with diamond history, emphasizing the discovery of synthetic diamond films and types of them. The second part (Chapters 4 to 6) concerns the beginning and new branches of the electrochemistry of diamond films. Chapters 7 and 8, which make up the third part of the book, summarize the studies of the diamond films on electroanalytical applications. The chapters focus on the use of these materials as sensors for detecting organic or inorganic species in order to propose the electroanalysis as an alternative quality control technique as well as for monitoring chemical species on air, soil, and water ecosystems. The fourth part (Chapters 9 to 19) is devoted to industrial applications of diamond films describing specific problems of particular interest for organic chemistry, energy conversion (fuel cells and batteries), and environmental aspects. Since diamond films are the most efficient materials for water treatment and water disinfection, Chapters 9 to 17 are devoted to these industrial applications in order to show the properties of these materials for environmental protection. The electrochemical oxidation with diamond films has been recognized as an electrochemical advanced oxidation process (EAOP), whereas the recent application of synthetic diamond films to other emerging EAOPs, such as electro-Fenton and photoelectro-Fenton, has also opened new prospects for wastewater remediation. The last part of the book (Chapters 20 to 21) focuses on innovative applications of diamond films, ranging from novel biosensors for chemical species of interest, as well as medical and clinical applications to neurochemistry and neuroscience by diamond microelectrodes and nanoelectrodes.

PREFACE

xxi

We strongly believe that this book will greatly promote research in this key field in the forthcoming years for the benefit of our society. Carlos A. Mart´ınez-Huitle Enric Brillas Editors

Preface to the Wiley Series on Electrocatalysis and Electrochemistry

This series covers recent advances in electrocatalysis and electrochemistry and depicts prospects for their contribution into the present and future of the industrial world. It aims to illustrate the transition of electrochemical sciences from its beginnings as a solid chapter of physical chemistry (covering mainly electron transfer reactions, concepts of electrode potentials and structure of electrical double layer) to the field in which electrochemical reactivity is shown as a unique chapter of heterogeneous catalysis, is supported by high-level theory, connects to other areas of science, and includes focus on electrode surface structure, reaction environment and interfacial spectroscopy. The scope of this series ranges from electrocatalysis (practice, theory, relevance to fuel cell science and technology) to electrochemical charge transfer reactions, biocatalysis and photoelectrochemistry. While individual volumes may appear quite diverse, the series promises updated and overall synergistic reports providing insights to help further our understanding of the properties of electrified solid/liquid systems. Readers of the series will also find strong reference to theoretical approaches for predicting electrocatalytic reactivity by such high-level theories as density functional theory. Beyond the theoretical perspective, further vehicles for growth are such significant topics such as energy storage, syntheses of catalytic materials via rational design, nanometer-scale technologies, prospects in electrosynthesis, new instrumentation and surface modifications. In this context, the reader will notice that new methods being developed for one field may be readily adapted for application in another. Electrochemistry and electrocatalysis have both benefited from numerous monographs and review articles due to their depth, complexity, and relevance to the practical world. The Wiley Series on Electrocatalysis and Electrochemistry is dedicated to present the current activity by focusing each volume on a specific topic that is timely and promising xxiii

xxiv

PREFACE TO THE WILEY SERIES ON ELECTROCATALYSIS AND ELECTROCHEMISTRY

in terms of its potential toward useful science and technology. The chapters in these volumes will also demonstrate the connection of electrochemistry to other disciplines beyond chemistry and chemical engineering, such as physics, quantum mechanics, surface science, and biology. The integral goal is to offer a broad-based analysis of the total development of the fields. The progress of the series will provide a global definition of what electrocatalysis and electrochemistry are now, and will contain projections about how these fields will further evolve in time. The purpose is twofold, to provide a modern reference for graduate instruction and for active researchers in the two disciplines, as well as to document that electrocatalysis and electrochemistry are dynamic fields that are expanding rapidly, and are likewise rapidly changing in their scientific profiles and potential. Creation of each volume required the editors involvement, vision, enthusiasm and time. The Series Editor thanks each Volume Editor who graciously accepted his invitation. Special thanks go to Ms. Anita Lekhwani, the Series Acquisitions Editor, who extended the invitation to edit this series to me and has been a wonderful help in its assembling process. Andrzej Wieckowski Series Editor

Contributors

Leonardo Santos Andrade, Departamento de Qu´ımica, Universidade Federal de Goi´as, Campus de Catal˜ao, Avenida Lamartine P. Avelar 1120, 75704-020 Catal˜ao, GO, Brazil John C. Angus, Department of Chemical Engineering, Case Western Reserve University, Cleveland, OH 44106-7217, USA Amara Apilux, Sensor Research Unit, Department of Chemistry, Faculty of Science, Chulalongkorn University, Patumwan, Bangkok 10330, Thailand Helmut Baltruschat, Institute for Physical and Theoretical Chemistry, Universit¨at Bonn, D 53117 Bonn, Germany ´ Erick Roberto Bandala Gonzalez, Universidad de las Am´ericas-Puebla, Departamento de Ingenier´ıa Civil y Ambiental Grupo de Investigaci´on en Energ´ıa y Ambiente, Sta. Catarina M´artir, Cholula-Puebla, M´exico J´essica Horacina Bezerra Rocha, Centro de Ciˆencias Exatas e da Terra, Departamento de Qu´ımica, Universidade Federal do Rio Grande do Norte, Campus Universit´ario-Lagoa Nova, CEP 59.072-970, Natal/RN, Brazil Karel Bouzek, Department of Inorganic Technology, Institute of Chemical Technology Prague, Prague, Czech Republic Enric Brillas, Laboratori d’Electroqu´ımica dels Materials i del Medi Ambient, Facultat de Qu´ımica, Departament de Qu´ımica F´ısica, Universitat de Barcelona, Mart´ı i Franqu`es 1-11 08028 Barcelona, Spain Erika Bustos, Centro de Investigaci´on y Desarrollo Tecnol´ogico en Electroqu´ımica, Parque Tecnol´ogico Quer´etaro, Sanfandila, C.P. 76703, Pedro Escobedo, Edo. de Quer´etaro, M´exico ˜ Pablo Canizares, Department of Chemical Engineering, Faculty of Chemical Sciences, Enrique Costa Building, Universidad de Castilla La Mancha, Campus Universitario s/n 13071 Ciudad Real, Spain xxv

xxvi

CONTRIBUTORS

Orawon Chailapakul, Department of Chemistry, Sensor Research Unit and Center for Petroleum, Petrochemicals, and Advanced Materials, Chulalongkorn University, Patumwan, Bangkok 10330, Thailand Pimkwan Chantarateepra, Program in Biotechnology, Faculty of Science, Chulalongkorn University, Patumwan, Bangkok 10330, Thailand Guohua Chen, Department of Chemical and Biomolecular Engineering, Hong Kong University Science & Technology, Clean Water Bay, Kowloon, Hong Kong, China Xueming Chen, Department of Environmental Engineering, Zhejiang University, 388 Yuhangtang Road, Hangzhou 310058, China Christos Comninellis, Ecole Polytechnique F´ed´erale, Lausanne, Group of Electrochemical Engineering, EPFL, 1015 Lausanne, Switzerland Yasuaki Einaga, Department of Chemistry, Keio University, 3-14-1 Hiyoshi, Yokohama 223-8522, Japan Orlando Fatibello-Filho, Departamento de Qu´ımica, Universidade Federal de S˜ao Carlos, C.P. 676, 13560-970 S˜ao Carlos-SP, Brazil Akira Fujishima,

President, Tokyo University of Science, Kagurazaka, Tokyo, Japan

Lu´ıs A. God´ınez, Centro de Investigaci´on y Desarrollo Tecnol´ogico en Electroqu´ımica, Parque Tecnol´ogico Quer´etaro, Sanfandila, C.P. 76703, Pedro Escobedo, Edo. de Quer´etaro, M´exico ´ Jos´e L. Guzman-Mar, Facultad de Ciencias Qu´ımicas, Centro de Laboratorios Especializados, Universidad de Nuevo Le´on, Pedro de Alba s/n, Cd. Universitaria, San Nicol´as de los Garza, NL, M´exico ´ırez, Facultad de Ciencias Qu´ımicas, Centro de Laborato´ Aracely Hernandez-Ram rios Especializados, Universidad de Nuevo Le´on, Pedro de Alba s/n, Cd. Universitaria, San Nicol´as de los Garza, NL, M´exico Lu´ıs Hinojosa-Reyes, Facultad de Ciencias Qu´ımicas, Centro de Laboratorios Especializados, Universidad de Nuevo Le´on, Pedro de Alba s/n, Cd. Universitaria, San Nicol´as de los Garza, NL, M´exico Katherine Holt, Department of Chemistry, University College London, Christopher Ingold Building, 20, Gordon St., London, WC1H 0AJ, UK Xin Jiang, Institute of Materials Engineering, University of Siegen, Paul-Bonatz-Str. 9-11, 57076 Siegen, Germany Agnieszka Kapalka, Ecole Polytechnique F´ed´erale, Lausanne, Group of Electrochemical Engineering, EPFL, 1015 Lausanne, Switzerland Axel Kirste, Kekul´e-Institut f¨ur Organische Chemie und Biochemie, Rheinische Friedrich-Wilhelms Universit¨at Bonn, Gerhard-Domagk-Str. 1, 53121 Bonn, Germany Carlos Alberto Mart´ınez-Huitle, Centro de Ciˆencias Exatas e da Terra, Departamento de Qu´ımica, Universidade Federal do Rio Grande do Norte, Campus Universit´ario-Lagoa Nova, CEP 59.072-970, Natal/RN, Brazil Yunny Meas, Centro de Investigaci´on y Desarrollo Tecnol´ogico en Electroqu´ımica, Parque Tecnol´ogico Quer´etaro, Sanfandila, C.P. 76703, Pedro Escobedo, Edo. de Quer´etaro, M´exico

CONTRIBUTORS

xxvii

Stamo Mentizi, Johannes Gutenberg-Universit¨at Mainz, Institut f¨ur Organische Chemie, Duesbergweg 10-14, 55128 Mainz, Germany Christoph E. Nebel, Fraunhofer-Institute for Applied Solid State Physics (IAF), Department Micro- and Nano-Sensors (GF5), Tullastrasse 72, 79108 Freiburg, Germany Marco Panizza, Department of Chemical and Process Engineering, Universit`a di Genova, P.le J.F. Kennedy 1, 16129 Genoa, Italy Bhavik A. Patel, Centre for Biomedical and Health Sciences Research, School of Pharmacy and Biomolecular Sciences, University of Brighton, Brighton, BN2 4GJ ´ Juan Manuel Peralta-Hernandez, Facultad de Ciencias Qu´ımicas, Centro de Laboratorios Especializados, Universidad de Nuevo Le´on, Pedro de Alba s/n, Cd. Universitaria, San Nicol´as de los Garza, NL, M´exico Yurnny V. Pleskov, Frumkin Institute of Physical Chemistry and Electrochemistry, Russian Academy of Sciences, Leninsky prospekt 31, 119991 Moscow, Russia Marco Antonio Quiroz-Alfaro, Universidad de las Am´ericas Puebla. Departamento de Ciencias Qu´ımico Biol´ogicas, Grupo de Investigaci´on en Energ´ıa y Ambiente, Sta. Catarina M´artir, Cholula-Puebla, M´exico Romeu Cardozo Rocha–Filho, Departamento de Qu´ımica, Universidade Federal de S˜ao Carlos, C.P. 676, 13560-970 S˜ao Carlos-SP, Brazil Manuel A. Rodrigo, Department of Chemical Engineering, Faculty of Chemical Sciences, Enrique Costa Building, Universidad de Castilla La Mancha, Campus Universitario s/n 13071 Ciudad Real, Spain ´ Cristina Saez, Department of Chemical Engineering, Faculty of Chemical Sciences, Enrique Costa Building, Universidad de Castilla La Mancha, Campus Universitario s/n 13071 Ciudad Real, Spain Giancarlo R. Salazar-Banda, Laborat´orio de Eletroqu´ımica e Nanotecnologia, Instituto de Tecnologia e Pesquisa, Universidade Tiradentes, Av. Murilo Dantas 300, Farolˆandia, 49032-490 Aracaju-SE, Brazil ´ Ana Sanchez-Carretero, Department of Chemical Engineering, Faculty of Chemical Sciences, Enrique Costa Building, Universidad de Castilla La Mancha, Campus Universitario s/n 13071 Ciudad Real, Spain Onofrio Scialdone, Dipartimento di Ingegneria Chimica, Gestionale, Meccanica e Informatica, Universit`a di Palermo, Viale delle Scienze, 90100, Palermo, Italy Virender K. Sharma, Florida Institute of Technology, 150 West University Boulevard Melbourne, Fl 32901, USA Weena Siangproh, Department of Chemistry, Faculty of Science, Srinakharinwirot University, Sukhumvit 23, Wattanna, Bangkok 10110, Thailand Ignasi Sir´es, Laboratori d’Electroqu´ımica dels Materials i del Medi Ambient, Facultat de Qu´ımica, Departament de Qu´ımica F´ısica, Universitat de Barcelona, Mart´ı i Franqu`es 1-11 08028 Barcelona, Spain Vadali V.S.S. Srikanth, School of Engineering Sciences and Technology, University of Hyderabad, Central University (P.O.), Hyderabad 500046, India

xxviii

CONTRIBUTORS

Siegfried R. Waldvogel, Institut f¨ur Organische Chemie, Johannes GutenbergUniversit¨at Mainz, Duesbergweg 10-14, 55128 Mainz, Germany Nianjun Yang, Fraunhofer-Institute for Applied Solid State Physics (IAF), Department Micro- and Nano-Sensors (GF5), Tullastrasse 72, 79108 Freiburg, Germany

Current density (mA cm–2)

0.20 (a) 0.15 0.10 (b) 0.05 (c) 0.00 0.2

0.0

0.4

0.6

0.8

Potential (V) vs. Ag/AgCl Figure 5.9 For caption see page 122.

counter electrode

BDD

stimulating electrode Reference electrode

anesthesia induction

Counter electrode

Stimulating electrode reference electrode

BDD

Brain Cortex Corpus Striatum Substantia Nigra Dopaminergic Neuron

Figure 5.14 For caption see page 127.

(c)

(b)

(d) 75 nA

4.0 3.0 2.0 1.0

0 nA

0.0

10 0 –10 0

20 40 60 Distance (µm)

80

Figure 6.3 For caption see page 145.

Current density(A m–2)

Anodic potential (V) vs. SCE

20

2.0

65

2.2

0.04 Thermodynamics

0.03

Pt

0.02 j (A cm–2)

Height (nm)

it/i(∞)

DDB

0.01 0

–0.01 –0.02

–2.0

–1.0

0.0

1.0

2.0

3.0

E (V) vs. NHE

–0.03 Figure 12.1 For caption see page 282.

90

2.5

155

3.5

195

4.1

Power supply V

A

catholyte

anolyte +

–

Electrochemical cell

Membrane Anode (BDD)

Cathode (AISI 304)

Out

Out

In

In

Figure 12.7 For caption see page 286.

(a)

(b)

(c)

(d)

µm 5

µm 5 4

4 3

3 2

2 1

1

Figure 16.16 For caption see page 398.

Figure 19.5 For caption see page 488.

Platelets Figure 20.6 For caption see page 523.

sensor

sensor

sensor

sensor

sensor BLOOD

Red blood cells

Activated platelets

Protein

Fibrin

Before stimulation

(a)

(b)

After stimulation at 10 Hz

Stimulator 210 µm

210 µm

Diamond Microelectrode (c) Imax = 10.4 pA

Stimulation Current (pA)

Diameter (µm) 30 µm

Figure 20.13 For caption see page 532.

Ileum–cross section

Submucosal plexus neurons EC Cells

Myenteric plexus neurons

Circular Muscle

Longitudinal Muscle

Enterochromaffin cells

Mucosal layer

SERT transporters present in enterocytes

Figure 20.14 For caption see page 533.

Myenteric plexus

(a)

Gastric Pit

(b) 8

Gastric glands Parietal cell

Current / nA

6

4

2 ECL-cell 0 0

100

200

300

500

400

Time (s)

Figure 20.15 For caption see page 534.

Figure 20.20 For caption see page 541. Au

Current (mA cm−2)

Pt Glassy Carbon

(H)SCD B:PCD (USU)

B:PCD (NRL)

B:(H)SCD –3

–2

–1

0

1

Potential (V) vs. SCE

Figure 21.1 For caption see page 552.

2

3

600

3

–4.2

pH2 = 1 bar 1 mbar 1 µbar

0 –1 –2

0

2

–3

CdS

2

–4.0

–4.4

µ

–χ

µe (eV)

1

EVAC

–1

–4.6 6 8 10 12 14 pH-Value

4

–2 –3

CdSe

µ

–5 –6 GaP

SiC

–7

GaAs Si

–4 –5

H-terminated Diamond

Ge

–6

Diamond

–7

Semiconductors Figure 21.3 For caption see page 554.

1500 Diamond 1000

Gold Silicon

500 Glassy Carbon

0

0

5

1 0

–4

Fluorescence intensity (a.u.)

Energy rel. to vacuum level (eV)

2

3

–3.8

Hydrogen Term Diamond EVBM

10

15

20

Hybridization cycles Figure 21.5 For caption see page 555.

25

30

(I)

(II)

Nanocrystalline Diamond

(a)

1 µm

1 µm

Silicon Polycrystalline Diamond

(b)

200 µm 200 µm Substrate Side

(c)

CL Intensity (a.u.)

4 x 106

FE(TO)

T = 16 K E = 13 kV I = 2 µA

3 x 106

2 x 106

1 x 106

0 5.0

0.5 nm

0.75

0.3 nm

0.50

0.0 nm

0.25

FE(LO) FE(TA) Γ FE(TO + O ) 5.1 5.2 5.3 5.4 Photon Energy (eV)

1.00

5.5

0

0.25

0.50

Figure 21.6 For caption see page 556.

0.75

0 1.00 µm

(a) Electrochemical nitrophenyl grafting

Nanotextures from Boron-Doped Diamond

(c) Probe DNA immobilization

Nanotextures from Boron-Doped Diamond

(b) Animation & cross-linker attachment

Nanotextures from Boron-Doped Diamond

(d) DNA hybridization & detection

Nanotextures from Boron-Doped Diamond

Figure 21.7 For caption see page 559.

500

400

[nm]

300

200

100

0 0

100

200

0.00

300

[nm]

400

500

0.50

[nm] Figure 21.8 For caption see page 560.

(a)

Current density (mA cm–2)

0.5 (b) (c) 0

–0.5

(d)

–0.5

0

0.5

Potential (V) vs. Ag/AgCl Figure 21.12 For caption see page 564.

1

(c)

(b)

(a)

(d)

Figure 21.13 For caption see page 565.

5

Sensor Area (norm.)

[×1013] 6 (a)

C –2 (F–2)

4

3 (c) 2

1

0 –0.5

2.2 2 1.8 1.6 1.4 1.2 1 0

(d)

5 10 15 Etching time (s)

20

(b)

0

0.5

1 1.5 2 E (V) vs. Ag/AgCl

2.5

Figure 21.16 For caption see page 568.

3

3.5

40

Height / A

32

24

16 0

1000

2000 Distance (nm)

Figure 21.21 For caption see page 573.

Figure 21.24 For caption see page 576.

3000

(

(a)

(

(b)

Figure 21.27 For caption see page 578.

Figure 21.28 For caption see page 579.

100 nN 120 nN >120 nN

6 Height (nm)

5 4 3

26 Å

8Å

2 1 0

0

2000 1000 Distance (nm)

3000

Figure 21.38 For caption see page 588.

Figure 21.41 For caption see page 590.

Figure 21.46 For caption see page 596.

Figure 21.47 For caption see page 597.

(a)

(b)

Figure 21.48 For caption see page 597.

(a) PEEK-Sample Holder (part 1)

(b) PEEK-Sample Holder (part 2, backside)

Figure 21.58 For caption see page 604.

(c) Mounted Sample (part 1 and 2)

A ds -D N

ds -D N

ds -D N

A

A

Pt

Debye Length λD DRAIN

Source

Figure 21.59 For caption see page 604.

(a) Single-stranded (ss) DNA

(b) Double-stranded (ds) DNA

Redox-Molecules diffuse into the ss-DNA film.

ss

Redox-Molecules are repelled by Coulomb Force

ds

Diamond

Diamond

Generation of Redox-Current

No Redox-Current

ss

4

5.0 x 10

0.0 4

–5.0 x 10

4

–1.0 x 10 –0.4 –0.2

0.0

0.2

0.4

Potential (V vs. Ag/AgCl)

0.6

Current density (A/cm2)

2

Current density (A/cm )

Ferrocyanide (5mM) redox reaction 1.0 x 104

1.0 x 104

ds

5.0 x 104 0.0 –5.0 x 104 4

–1.0 x 10 –0.4

–0.2 0.0 0.2 0.4 0.6 Potential (V vs. Ag/AgCl)

Figure 21.62 For caption see page 607.

Part I

Synthesis of Diamond Films

1 Electrochemistry on Diamond: History and Current Status John C. Angus

1.1

ENABLING TECHNOLOGIES

The use of diamond in electrochemistry was not anticipated during the initial explosion of interest in low-pressure diamond growth during the 1980s. Most attention was focused on electronic, optical, and mechanical applications (e.g., for high temperature electronic devices, radiation detectors, high voltage switches, x-ray windows, audio speaker diaphragms, and protective and tool coatings). Little attention was paid to electrochemical applications. In this chapter, we describe the enabling technologies underlying diamond electrochemistry, the early work on diamond electrochemistry itself, and two major threads of current research. The present status of the field is covered in the other chapters of this volume. 1.1.1

Chemical Vapor Deposition of Diamond

Diamond electrochemistry has been made possible by development of methods for the chemical vapor deposition of diamond. The first reported growth of diamond at sub-atmospheric pressure, where diamond is metastable with respect to graphite, was by Eversole [1] at Union Carbide Corporation. He grew diamond on high-surface area powder from methane and carbon monoxide. This effort was ultimately abandoned after Synthetic Diamond Films: Preparation, Electrochemistry, Characterization, and Applications, First Edition. Edited by Enric Brillas and Carlos Alberto Mart´ınez-Huitle. © 2011 John Wiley & Sons, Inc. Published 2011 by John Wiley & Sons, Inc.

3

4

ELECTROCHEMISTRY ON DIAMOND: HISTORY AND CURRENT STATUS

the announcement of the successful growth of diamond at high pressures by General Electric Corporation [2]. Eversole’s work was later extended by Angus et al. at Case Western Reserve University [3–7] and by Deryagin and a large group at the Physical Chemistry Institute in Moscow [8–13]. The Angus group showed the beneficial effect of added hydrogen on diamond yield, developed methods for removing unwanted graphitic carbon using atomic hydrogen, and grew boron-doped diamond by chemical vapor deposition [5]. In their first studies, the Soviet workers reported the growth of filamentary diamond whiskers using a vapor-liquid-solid (VLS) technique with molten iron and nickel [8,9]. Surprisingly, this process has received little attention in subsequent years. From the same group at the Physical Chemistry Institute, Varnin et al. [11] grew diamond thin films from the gas phase; Spitsyn, Bouilov, and Deryagin [13] reported both diamond films and isolated diamond crystals grown on copper substrates by chemical vapor transport from a graphite source. These very early studies were the first to show that diamond could be grown at conditions in which it is not the stable phase. At the time, diamond synthesis at low pressures was widely believed to be impossible and in violation of the second law of thermodynamics. Although the early studies showed the feasibility of diamond synthesis by chemical vapor deposition, the growth rates were low. This situation changed dramatically in the mid-1980s. A group at the National Institute for Research in Materials in Tsukuba, Japan, under the leadership of Nobuo Setaka achieved large increases in growth rates by activating the gas phase during growth with a hot filament [14,15], a microwave discharge [16], and an RF discharge [17,18]. Matsui, Matsumoto, and Setaka [19] also performed careful characterization of their product diamond. Since the mid-1980s, the interest in chemical vapor deposition of diamond has expanded enormously, and it is now used to grow extremely high-quality diamond in many shapes and sizes. 1.1.2

Doping of Diamond

For most electrochemical applications, it is necessary to use conducting diamond obtained by doping with the group III element, boron. The effect of boron on the electrical conductivity of diamond has been well studied and is covered in several classic review volumes [20–24]. Poferl, Gardner, and Angus [5] in 1973 were the first to grow boron-doped conducting diamond by chemical vapor deposition. This study was done to provide further confirmation of metastable diamond growth. Since then, numerous studies of diamond doping have been made. The boron-doping agent is usually added through small amounts of diborane, trimethyl boron, or organic borates in the source gases, but solid-state boron sources have also been successfully used. Reviews by Pan and Kania [25] and Werner and Locher [26] describe the state of the art in the mid-1990s; more recent discussions [27–29] are available. At low concentrations, boron promotes p-type semiconductivity in diamond with an acceptor level 0.37 eV above the valence band [30]. At boron concentrations higher than 1017 to 1018 cm−3 , an impurity band forms and the acceptor gap is reduced [31–34]. At boron concentrations from 1020 to 1021 cm−3 , diamond is a semimetal, with resistivities as low as 0.001 -cm. For electrosynthesis and electrodestruction applications, high conductivities are desirable; for photoelectrochemistry and semiconductor electrochemistry, low conductivity electrodes are used. The uniformity of boron concentration is an important variable in electrochemical applications. Spitsyn et al. in 1981 [13] and Janssen et al. in 1990 [35] showed that more boron was incorporated in (111) growth sectors than in (100) sectors.

1.2 FIRST STUDIES OF THE ELECTROCHEMISTRY ON DIAMOND

1.1.3

5

Surface Characterization of Diamond

Diamond, graphene and carbon nanotubes are unique among the common semiconductors in having no stable solid surface oxide. Furthermore, the chemistry of the diamond surface is closely related to well-known organic chemical processes. As a consequence, it is possible to tailor the properties of the diamond surface for sensor and other applications by changing the surface functionalization. The most significant early paper on diamond surfaces was the LEED study by Lander and Morrison at Bell Laboratories in 1966 [36]. They showed that hydrogen-termination of (111) diamond surfaces eliminated the reconstruction of the clean surface and gave an (essentially) bulk-terminated surface. They also found significant atom mobility on the diamond surface between 900◦ C and 1400◦ C. Lander and Morrison explicitly stated that this behavior could make epitaxial extension of diamond feasible. Another important early study of diamond surfaces and their interaction with gases was by Lurie and Wilson in 1977 [37].

1.2

FIRST STUDIES OF THE ELECTROCHEMISTRY ON DIAMOND

Even after the chemical vapor deposition of diamond became more or less routine, there was little availability of boron-doped conducting diamond. Therefore, some of the earliest electrochemical studies done on diamond were performed on samples in which conductivity was induced by damage from ion implantation. In 1983, Iwaki et al. [38] of the Institute of Physical and Chemical Research in Hirosawa, Japan, reported the use of diamond electrodes made conductive by ion implantation of nitrogen, argon, and zinc. However, the ion implantation converted the diamond surface into nondiamond carbon [38,39] so the measurements did not reflect the electrochemical characteristics of diamond itself. 1.2.1

From 1987 to 1996

The first measurements of the inherent electrochemical properties of diamond were by Yuri Pleskov and co-workers at the Frumkin Institute of Electrochemistry in Moscow. In 1987, they reported the photoelectrochemical properties of diamond electrodes [40]. They found a photoresponse of diamond at sub-bandgap wavelengths that they attributed to excitation of electrons from mid-gap defect states to the conduction band. They also observed sluggish evolution of hydrogen at cathodic potentials. Subsequently, in 1989 Natishan and Morrish at the Naval Research Laboratory in Washington, D.C., described the electrochemical properties of composite diamond/molybdenum electrodes in which the diamond acted as an inert insulating material [41]. During the early 1990s, Pleskov and co-workers continued and expanded their earlier work on diamond electrochemistry. They analyzed the impedance of conducting polycrystalline diamond thin films between two ohmic contacts and showed that intercrystalline boundaries are capacitive/resistive barriers in the bulk of diamond [42,43]. This same group continued its photoelectrochemical studies [44,45] and measured the minority carrier diffusion length and the acceptor concentrations in diamond [46]. A large group at the University of Tokyo led by Akira Fujishima made extensive progress on diamond electrochemistry during the early 1990s. Patel, Hashimoto, and

6

ELECTROCHEMISTRY ON DIAMOND: HISTORY AND CURRENT STATUS

Fujishima examined the photoresponse of diamond [47,48]. They showed that the diamond electrodes acted as a p-type semiconductor; from the flat band potential, they deduced that the bottom of the conduction band of diamond was near the vacuum level. Tenne and others of this same group exploited the large overpotentials for hydrogen evolution for the reduction of nitrate to ammonia on boron-doped diamond electrodes [49,50]. In 1993 Marchywka et al. [51] described an interesting process in which they used low-energy carbon ion implantation, followed by electrochemical etching of the damaged region, to separate thin diamond layers from bulk diamond crystals. Also in 1993, Ramesham and co-workers [52] deposited diamond films on both glassy carbon and graphite to make a composite electrode that they said “may have some use in electroanalysis”. Miller et al. in 1994 used ion implantation of 140 keV Co ions to form a patterned, nondiamond conductive region on a planar diamond electrode surface [53]. In the early and mid-1990s, Swain and his colleagues published a series of papers [54–58] that demonstrated many of the essential properties of diamond electrodes: their low capacitance, featureless background current, large signal to noise ratio, and chemical stability. They suggested that these properties made diamond electrodes suitable for electroanalysis and sensors. In 1995 a group of researchers at the Shanghai Institute of Metals also reported on the stability and reproducibility of diamond electrodes and noted their potential for sensor applications [59–61]. The large overpotential for hydrogen evolution on diamond was observed in the earliest work [49,50], but the large overpotential for oxygen evolution did not become apparent until high-quality diamond electrodes, without significant nondiamond carbon, were used. In 1995 and 1996, Martin et al. [62], Argoitia et al. [63], and Martin et al. [64] demonstrated that high-quality diamond electrodes did not evolve hydrogen until −1.25 V and oxygen until +2.3 V (versus the standard hydrogen electrode). Bouamrane et al. [65] reported a similar observation in 1996. The high overpotentials for both oxygen and hydrogen evolution on diamond give rise to an extremely wide potential window of water stability that is found on no other electrode material. Figure 1.1 shows the current-voltage characteristics for water electrolysis on high- and low-quality diamond, glassy carbon, and platinum [63]. The wide potential window and the lack of background current within the potential window for high-quality diamond are apparent. 1.2.2

From 1996 to Present

In the latter part of the 1990s, several intensive efforts were made to understand the nature of charge transfer at diamond electrodes and how it differed from conventional carbon electrodes. In 1996, Vinokur et al. [66] performed a careful analysis of electron transfer to both lightly and heavily doped diamond electrodes. Highly reversible electrode behavior was found for heavily boron-doped electrodes, especially at potentials more positive than −0.50 V versus the saturated calomel electrode. Lightly doped, semiconducting electrodes showed very slow kinetics at more negative potentials, which was attributed to a lack of available states in the bandgap of diamond. During this same time period, the Swain [67] group measured the boron distribution on diamond electrodes and Granger et al. [68] presented a detailed study of the electrode kinetics of diamond from two different sources. They examined the difference in electrode kinetics for processes involving outer sphere

1.2 FIRST STUDIES OF THE ELECTROCHEMISTRY ON DIAMOND

(a) Platinum

(b) High-Quality Diamond 30

15 0 −1

−15

1

2

3

i (mA / cm−2)

i (mA / cm−2)

30

−2

15 0 −2

−1 −15

V vs. SHE(V)

1

2

3

i (mA / cm−2)

i (mA / cm−2)

0

V vs. SHE(V)

1

2

3

30

15

−30

3

(d) Glassy Carbon

30

−15

2

V vs. SHE(V)

(c) Low-Quality Diamond

−1

1

−30

−30

−2

7

15 0 −2

−1 −15 −30 V vs. SHE(V)

Figure 1.1 Early comparison of voltammograms of water electrolysis on (a) platinum, (b) high-quality diamond, (c) low-quality diamond, and (d) glassy carbon. The electrolyte was 0.50 M H2 SO4 ; potentials are measured versus the standard hydrogen electrode. The wide potential window between hydrogen evolution at ∼ − 1.5 V and oxygen evolution at ∼ + 2.5 V and lack of background current for the high-quality diamond electrode are evident. Reprinted from [63].

electron transfer and for reactions that go through adsorbed intermediates; they also studied the effect of hydrogen-termination and surface oxidation on the electrode kinetics. In 1999 Martin and associates [69] showed that increasing amounts of sp2 , nondiamond carbon on the diamond surface reduced the potential window of water stability and made the electrodes behave more like glassy carbon or pyrolytic graphite. The results of these and other studies showed the intrinsic differences between diamond and conventional carbon electrodes. Many of these differences arise from the very different types of surface structure and functionalities that are found on diamond electrode surfaces. One of the most striking features is the relatively inert, hydrogen-terminated diamond surface. The high overpotentials for hydrogen and oxygen evolution (and other reactions) are attributed, at least in part, to the limited adsorption of intermediate species on diamond. High overpotentials are not observed for rapid outer-sphere reactions that do not rely on an adsorbed intermediate. However, it should be emphasized that the diamond surface is not completely inert. Anodic polarization and liberation of oxygen changes the surface properties of a previously hydrogen-terminated surface [40,64,69–71]. These changes have been attributed to the replacement of hydrogen with oxygen functionalities to the surface [72]. Rao et al. [72] also showed that changing the surface functionalization changed the absolute position of the band edges and hence the flat band potential. Furthermore, there is evidence that in some cases oxidized diamond surfaces have a greater potential window of water stability than do hydrogenated surfaces [73–76].

8

ELECTROCHEMISTRY ON DIAMOND: HISTORY AND CURRENT STATUS

Of great interest is the possibility that the chemically bound hydrogen participates in the hydrogen evolution reaction generating a free radical site, for example, CH + H+ + e− → C • + H2

(1.1)

Anderson and Kang [77] showed that the energetics of this reaction are favorable. Furthermore, Yagi et al. [78] demonstrated deuterium exchange between the solution and a hydrogenated diamond surface during hydrogen evolution. Argoitia et al. [63] and Martin et al. [69] reported a rather surprising result that oxygen was added to hydrogenterminated diamond surfaces during liberation of hydrogen. The preceding results show that the electrochemical processes on diamond are significantly more complex than if the diamond were a completely inert source or sink of electrons. The early work on the electrochemistry of diamond was summarized in reviews by Tenne and Levy-Clement [79], Swain, Anderson, and Angus [80], Angus et al. [81], Kobashi [82], and Pleskov [83].

1.3 DEVELOPMENT OF ELECTROCHEMICAL APPLICATIONS OF DIAMOND By 2000 boron-doped diamond electrodes were receiving much attention. In addition to the very low background currents, low capacitance, and extremely wide potential window for water stability, diamond electrodes were shown to be extremely stable in many corrosive environments. Furthermore, the electrodes resisted fouling, especially when hydrogen-terminated. These properties began to be exploited in a number of laboratories around the world for a variety of applications. 1.3.1

Surface Functionalization

A natural extension of the early surface studies has been the extensive efforts on adding specific functional groups to the diamond surface. One of the first reported studies was that of Freedman and Stinespring [84] who in 1990 fluorinated diamond surfaces with an atomic beam of fluorine atoms. Smetkowski and Yates [85] in 1996 added fluorine to diamond surfaces by x-ray irradiation of perfluoroalkyl iodide layers on the surface. Ohtani et al. [86] in 1998 treated diamond with chlorine under ultraviolet irradiation; subsequent heating in boiling pyridine added quaternary pyridinium salts to the diamond. In 1999 Kuo, McCreery, and Swain [87] electrochemically coupled phenyl groups to diamond surfaces using phenyl diazonium. Polarization-induced modifications of diamond surfaces were described by Goeting et al. [88] in 1999. Throughout the early 2000s, there were many studies of addition of hydrocarbon functional groups to diamond. In 2001 Mathieu [89] reported functionalizing diamondcovered silicon wafers to improve biocompatibility. Wang et al. [90] used Diels-Alder chemistry to functionalize diamond with hydrocarbons. The theory of Diels-Alder reactions on diamond surfaces was discussed by Fitzgerald and Doren [91] and Lu et al. [92]. The reactivity of the 2 × 1 reconstructed (100) surface of diamond to 1, 3 butadiene and cyclopentene was reported by Russell et al. [93]. Hovis et al. [94] and Strother et al. [95] photochemically functionalized polycrystalline diamond surfaces with several

1.3 DEVELOPMENT OF ELECTROCHEMICAL APPLICATIONS OF DIAMOND

9

organic molecules. Another approach was reported by Ida et al. [96,97], who treated hydrogenated diamond surfaces with organic peroxides to add carboxylic acids and other functionalities. In 2002 Bent [98] reviewed the status of chemical functionalization of group IV semiconductor surfaces, including diamond. 1.3.2

Destruction of Wastes

One early application of conducting diamond was the anodic destruction of organic wastes. This process takes advantage of the very high anodic overpotentials that can be achieved on diamond electrodes as well as their chemical and dimensional stability under extreme conditions. Carey, Christ, and Lowery [99] used this method to oxidize photo solutions, phenol, hydroquinone, sodium salts of EDTA, and formic, oxalic, and malonic acids. From 1999 to 2001, Comninellis and co-workers [100–110] published an extensive series of papers on the use of boron-doped diamond electrodes for destruction of wastes. They showed that highly oxidizing species, for example, the persulfate ion, could be generated on diamond electrodes at anodic potentials well above the reversible potential for oxygen evolution [104]. They also demonstrated that hydroxyl groups generated at high anodic potentials were efficient oxidizing agents that kept organic films from forming on the diamond anodes [102,103,107,110]. Other workers used diamond anodes to both detect and destroy organic materials in waste streams. Davila-Jimenez et al. [111] reported in 2000 that textile dyes could be anodically treated using diamond electrodes. In 2002, Terashima et al. [112] used diamond electrodes to detect chlorophenols in drain waters. Bellagamba et al. [113] employed diamond anodes to oxidize dissolved polyacrylates in water, and Van Hege, Verhaege, and Verstraete [114] at the State University of Ghent oxidized organic components and ammonia in the concentrated filtrate from reverse osmosis membranes on diamond anodes. Lawrence et al. [115] detected and oxidized sulfide ion on boron-doped diamond electrodes; Van Andel and Janssen [116] described a process for the regeneration of chrome solutions using diamond electrodes, and Canizares et al. [117,118] described the destruction of phenol wastes using diamond anodes. Another interesting application is the use of boron-doped diamond anodes for the highly oxidizing electro-Fenton reaction [119,120]. Electrochemical oxidation on diamond electrodes is an active and ongoing field; recent work is described in later chapters of this book. 1.3.3

Sensors and Electroanalysis

The potential of diamond electrodes for sensor and analytical applications was recognized early by a number of workers: Ramesham et al. [52], Swain and Ramesham [54], Awada, Strojek, and Swain [58], and Zhu et al. [59]. The first applications followed shortly after this recognition. In 1996, Argoitia et al. [63] reported that the cerium/cerous redox couple could be detected on diamond. Strojek et al. [121] found enhanced signal to noise ratios at diamond electrodes. Also in 1996, Davidson et al. at Vanderbilt University developed diamond-based hydrogen [122] and oxygen [123] gas sensors. Use of diamond electrodes for analysis of biological molecules was an early research direction. In 1998 a group at Heriot Watt University in Edinburgh described the use of heavily boron-doped diamond electrodes for the determination of glucose [124]. In 1999 Rao et al. [125] from the University of Tokyo reported the electrochemical oxidation of nicotinamide adenine dinucleotide (NADH) and sensors based on enzyme-catalyzed

10

ELECTROCHEMISTRY ON DIAMOND: HISTORY AND CURRENT STATUS

reactions involving NADH as a cofactor. Methods for determining sulfadiazine [126] as well as histamine and serotonin [127] were developed by Sarada et al., and Popa, Tryk, and Fujishima [128] developed methods for purines. Fujishima, Rao, and Tryk [129] demonstrated that diamond electrodes could be used for the detection of trace amounts of lead, without the use of mercury. In 2001 two reports described the use of transparent diamond electrodes for making in situ spectroscopic observations. In one, the addition of hydroxyl groups to diamond was observed [130] and in the other the electro-oxidation of ferrocyanide and electroreduction of methyl viologen was monitored [131]. Also in 2001 Hayashi et al. [132] used a double layer of undoped- and boron-doped diamond between two platinum electrodes to detect phosphine, diborane, and arsine. Kawarada et al. [133] fabricated electrolytesolution-gate field effect transistors (FETs) as biocompatible ion sensors. Kohn et al. [134] discussed the prospects of diamond micro devices and reviewed the use of diamond in electromechanical, electrothermal, and electrochemical sensors [135]. Saterlay, Foord, and Compton [136] developed a diamond sensor to be used with ultrasound for the detection of 4-chlorophenol in aqueous solution. Fujishima and Rao [137] described diamond detectors for flow-injection analysis of chlorophenols and theophylline. Development of diamond-based sensors is continuing in many laboratories around the world. Carbon-based sensors, including diamond, have been the subjects of an extensive review by McCreery [138]; other recent reviews on diamond-based sensors are by Chailapakul, Siangproh, and Tryk [139] and Zhou and Zhi [140].

1.4

OTHER DIRECTIONS

Two current directions of research on diamond electrochemistry deserve mention: bioelectronic applications of diamond and the anomalous surface conductivity of diamond. Both areas are the subject of much current research and both appear to have significant long-range implications. 1.4.1

Biolectronic Applications

Diamond as a platform for bioelectronic and bioelectrochemical applications is becoming a major research tool. These applications exploit the biocompatibility of diamond, its chemical inertness, and the ability to tailor its surface chemistry to permit attachment of biologically active molecules (e.g., DNA) for bioelectronics and biosensing. Diamond has significant advantages over silicon and the other common semiconductors because it does not have a solid surface oxide interposed between it and the external environment. Biosensor work using diamond has been a continuing focus. These studies include not only early sensor studies mentioned in the previous section but also the electrochemical analysis of nucleic acids [141] and the determination of single and double-stranded DNA [142] at boron-doped diamond electrodes. Another direction was explored by Tachiki et al. [143] who measured the inter-relationship between the surface charging of diamond and the physisorption of DNA. In 2001 Adamschik et al. described a diamond lab on a chip for biochemical applications including DNA synthesis [144]. In 2002 Ushizawa et al. immobilized DNA on diamond powder through an ester linkage [145]. The covalent bonding of the DNA to the diamond was verified by diffuse reflectance infrared spectroscopy. In 2002 and 2003, Hamers and co-workers

1.4 OTHER DIRECTIONS

11

used a photochemical scheme to obtain DNA functionalization of nanocrystalline and polycrystalline diamond films through amine groups [146,147]. In 2003 Takahashi et al. [148] sequentially hydrogenated, chlorinated, aminated, and finally carboxylated a diamond substrate to provide a platform for chemically binding DNA. Also in 2003, Wenmackers et al. [149] used thymidine as a linker molecule to immobilize DNA on diamond; the bound DNA was confirmed by confocal fluorescence microscopy. Functionalization of diamond with DNA for sensor and other applications remains of major current research interest. A significant body of work has been achieved by Hamers and co-workers at the University of Wisconsin. These studies include measuring changes in surface electrical properties of DNA modified diamond films as a function of exposure to complementary and noncomplementary DNA molecules [150], the fabrication of a biologically sensitive field effect transistor on a nanocrystalline diamond film [151], and an electrically addressable functionalized film [152]. This same group also compared the thermal stability of DNA-modified diamond and silicon surfaces [153] and developed methods for functionalizing diamond surfaces to resist protein adsorption, while still permitting specific binding for biorecognition [154,155]. Nebel and co-workers [156] at the Diamond Research Center in Tsukuba, Japan, studied the geometric orientation of DNA covalently bonded to single crystal diamond and the photochemical attachment of oriented amine linkers to hydrogen-terminated single crystal diamond [157]. This same group used periodically arranged benzene linker molecules, also on single crystal diamond, to obtain dense DNA films [158]. Recent reviews on bioelectronics and bioelectrochemistry have been written by Linares, Doering, and Linares [159], Nebel et al. [160] and Vermeeren et al. [161]. 1.4.2

Anomalous Surface Conductivity of Diamond

Most of the interaction between diamond science and electrochemistry exploits the unusual properties of diamond for electrochemical applications. There is one important example of the reverse in which electrochemical principles have lead to a deeper understanding of the properties of diamond. In 1989 Landstrass and Ravi [162] reported a curious observation that hydrogenterminated diamond developed a pronounced p-type surface conductivity upon exposure to air. This observation was extremely unusual because surface conductivity had not been reported previously despite more than a century of study of the electronic properties of diamond. Subsequent research, initially by Gi et al. [163–165], started to unravel the cause of the conductivity. Gi showed that the conductivity increased with exposure to acidic gases such as NO2 , decreased in the presence of basic gases such as NH3 [163,164], and that the sheet concentration of holes was as high as 7 × 1013 cm−2 [164]. In order to explain the effect, Gi proposed that the holes arose from the oxidation of the terminal hydrogen atoms by the hydronium ion, H3 O+ [165]. Gi’s mechanism was one of several put forward to explain the surface conductivity, including the presence of acceptor sites arising from subsurface hydrogen. Maier et al. [166,167] framed the surface oxidation process in terms of an electrochemical redox reaction. They proposed that the hydrogen redox couple, 2H+ + 2e− = H2 , in an adsorbed water film acted as an external electrochemical acceptor to the electrons in the diamond. The electron chemical potential (Fermi energy) of the redox couple lay below the Fermi energy of hydrogen-terminated diamond, which permitted the couple to act as an external electron acceptor, for example, to oxidize the diamond. In this model, the

12

ELECTROCHEMISTRY ON DIAMOND: HISTORY AND CURRENT STATUS

role of the hydrogen-termination was to provide a surface dipole that raised the electronic band structure of diamond so that the electron transfer could take place. This proposal also received rather limited support, in part because the partial pressure of H2 in the atmosphere is so low that it was difficult to understand how it could support a reversible redox couple. Subsequently, Foord et al. [168] and Chakrapani et al. [169] gave evidence that the electrochemical mechanism was correct, but that the oxygen redox couple, 4H+ + 4e− + O2 = 2H2 O, was responsible for the effect. Calculations of Petrini and Larson [170] showed that the oxygen and ozone redox couples were both more favorable for electron transfer from diamond than the hydrogen couple. Chakrapani et al. [171] used macroscopic size samples of diamond and water to demonstrate that the changes in pH and dissolved oxygen concentration were consistent with electron transfer to the oxygen redox couple. It is of considerable interest that electron transfer to an external redox couple was proposed by Shapoval et al. earlier (in 1995) to explain conductivity changes in diamond exposed to molten salts [172]. The electrochemical transfer doping of diamond is still a subject of active research. Much more remains to be done to achieve a full understanding of the effect and to explore potential applications. It should be noted in this context that the surface conductivity has been used as the basis of interesting field effect transistors, most notably by Kawarada et al. [133] and Denisenko, Aleksov, and Kohn [173]. The understanding that an electrochemical mechanism played a role in mediating the electrical properties of diamond led to attempts to find the effect in other systems. Ristein [174] discussed the general nature of transfer doping as an alternative to conventional impurity doping. Strobel et al. [175] demonstrated transfer doping of diamond to an adsorbed surface layer of fullerene molecules, C60 . Chakrapani et al. reported electrochemical transfer doping of carbon nanotubes [176] and that the photoluminescence of GaN and ZnO could be mediated by electrochemical transfer doping by changing the

300

Number of publications

250 200 150 100 50 0 1990

1995

2000 Publication year

2005

2010

Figure 1.2 Number of publications on diamond electrochemistry versus year of publication. Three publications appeared prior to 1990: Iwaki et al. [38] in 1983 and Pleskov et al. [40] and Natishan and Morrish [41] in 1989. Source: ISI Web of Knowledge; the count was limited to papers identified by the keyword combinations (diamond electrochemistry), (diamond electrochemical) etc.

REFERENCES

13

acidity or basicity of the ambient humid air [177]. A very interesting development is the prediction that electrochemical charge transfer to the oxygen redox couple can be used to alter the conductivity of graphene [178]. Chen et al. [179] have recently published a comprehensive review of surface transfer doping.

1.5

CONCLUSIONS

The current status of diamond electrochemistry can be found in a recent review volume [180] and in the following chapters of this volume. Major areas of expanding interest include waste treatment using diamond anodes, electroanalysis on diamond electrodes, diamond-based electrochemical sensors, diamond electrodes for functional stimulation of muscle, diamond as a substrate for affixing DNA and other molecules of biological interest, and electronic devices based on diamond surface conductivity. The interest in these and other types of electrochemistry involving diamond is continuing to increase. The number of published papers concerned with diamond electrochemistry as a function of time is shown in Figure 1.2.

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

W.G. Eversole, US patents 3,030,187 and 3,030,188, 1962. F.P. Bundy, H.T. Hall, H.M. Strong, R.H. Wentorf, Nature 1955, 176 , 51–55. J.C. Angus, H.A. Will, W.S. Stanko, J. Appl. Phys. 1968, 39 , 2915–2922. J.C. Angus, D.J. Poferl, N.C. Gardner, S. Chauhan, T.J. Dyble, P. Sung, Sint. Almazy 1971, 3 , 30–31 (in Russian). D.J. Poferl, N.C. Gardner, J.C. Angus, J. Appl. Phys. 1973, 44 , 1428–1434. S.P. Chauhan, J.C. Angus, N.C. Gardner, J. Vac. Sci. Tech. 1974, 11 , 423. S.P. Chauhan, J.C. Angus, N.C. Gardner, J. Appl. Phys. 1976, 47 , 4746–4754. B.V. Deryagin, D.V. Fedoseev, V.M. Lukyanovich, B.V. Spitsyn, V.A. Ryabov, A.V. Lavrentiev, J. Cryst. Growth 1968, 2 , 380–384. B.V. Deryagin, D.V. Fedoseev, Russ. Chem Rev . 1970, 39 , 783–788. D.V. Fedoseev, K.S. Uspenskaya, V.P. Varnin, S.P. Vnukov, Izv. Akad. Nauk SSSR, Seriya Khimicheskaya 1978, 6 , 1252–1256. V.P. Varnin, B.V. Deryagin, D.V. Fedoseev, I.G. Teremetskaya, A.N. Khodan, Kristallografiya 1977, 22 , 893–896. D.V. Fedoseev, B.V. Deryagin, V.P. Varnin, G.K. Sokolina, Doklady Akad. Nauk SSSR 1979, 247 , 1201–1204. B.V. Spitsyn, L.L. Bouilov, B.V. Deryagin, J. Cryst. Growth 1981, 52 , 219–226. S. Matsumoto, Y. Sato, M. Kamo, N. Setaka, Jpn. J. Appl. Phys. 1982, 21 , part 2 , L183–L185. S. Matsumoto, Y. Sato, M. Tsutsumi, N. Setaka, J. Mater. Sci . 1982, 17 , 3106–3112. M. Kamo, Y. Sato, S. Matsumoto, N. Setaka, J. Cryst. Growth 1983, 62 , 642–644. S. Matsumoto, J. Mater. Sci. Lett . 1985, 4 , 600–602. S. Matsumoto, M. Hino, T. Kobayashi, Appl. Phys. Lett . 1987, 51 , 737–739. Y. Matsui, S. Matsumoto, N. Setaka, J. Mater. Sci. Lett . 1983, 2 , 532–534. R. Berman, Physical Properties of Diamond , Clarendon Press, Oxford, 1965.

14

ELECTROCHEMISTRY ON DIAMOND: HISTORY AND CURRENT STATUS

21. J. E. Field, ed., The Properties of Diamond , Academic Press, London, 1979. 22. G. Davies, Diamond , Adam Hilger Ltd., Bristol, 1984. 23. J. Wilks, E. Wilks, Properties and Applications of Diamond , Butterworth Heinemann, Oxford, 1991. 24. J.E. Field, ed., Properties of Natural and Synthetic Diamond , Academic Press, London, 1992. 25. L.S. Pan, D.R. Kania, eds, Diamond: Electronic Properties and Applications, Kluwer Academic Publishers, Boston, 1995. 26. M. Werner, R. Locher, Rep. Prog. Phys. 1998, 61 , 1665–1710. 27. C.E. Nebel, J. Ristein, eds., Thin Film Diamond I, Semiconductors and Semimetals, Volume 76 , Elsevier Academic Press, Amsterdam, 2003. 28. C.E. Nebel, J. Ristein, Eds., Thin Film Diamond II, Semiconductors and Semimetals, Volume 77 , Elsevier Academic Press, Amsterdam, 2004. 29. R. S. Sussman, ed., CVD Diamond for Electronic Devices and Sensors, John Wiley & Sons, Ltd, Chichester, 2009. 30. A.T. Collins, Phil. Trans. R. Soc. A 1993, 342 , 233–234. 31. D. Narducci, C.R. Guarnieri, J.J. Cuomo, J. Electrochem. Soc. 1991, 138 , 2446–2451. 32. Y.F. Tsai, K.P. Ananthanarayanan, P.J. Gielisse, S.S. Mitra, J. Appl. Phys. 1972, 43 , 3677–3682. 33. J. Mort, K. Okumura, M.A. Machonkin, Phil. Mag. B 1991, 63 , 1031–1036. 34. K. Nishimura, K. Das, J.T. Glass, J. Appl. Phys. 1991, 69 , 3142–3148. 35. G. Janssen, W.J.P. van Enckevort, W. Vollenberg, L.J. Giling, Diamond Relat. Mat . 1990, 1 , 789–800. 36. J. J. Lander, J. Morrison, Surf. Sci . 1966, 4 , 241–246. 37. P. G. Lurie, J. M. Wilson, Surf. Sci . 1977, 65 , 453–475. 38. M. Iwaki, S. Sato, K. Takahashi, H. Sakairi, Nucl. Instrum. Methods Phys. Res. 1983, 209–210 , 1129–1133. 39. M. Iwaki, K. Takahashi, H. Watanabe, S. Sato, RIKEN Review 1996, 12 , 27–28. 40. Y. V. Pleskov, A. Y. Sakharova, M. D. Krotova, L. L. Bouilov, B. V. Spitsyn, J. Electroanal. Chem. 1987, 228 , 19–27. 41. P.M. Natishan, A. Morrish, Mat. Lett . 1989, 8 , 269–272. 42. A.Ya. Sakharova, A.E. Sevastyanov, Y.V. Pleskov, G.L. Teplitskaya, V.V. Surikov, A.A. Voloshin, Soviet Electrochemistry 1991, 27 , 239–244. 43. A. Sakharova, L. Nykos, Y.V. Pleskov, Electrochim. Acta 1992, 37 , 973–978. 44. Y.V. Pleskov, A.Y. Sakharova, E.V. Kasatkin, VA. Shepelin, J. Electroanal. Chem. 1993, 344 , 401–404. 45. A.Y. Sakharova, Y.V. Pleskov, F. Diquarto, S. Piazza, C. Sunseri, I.G. Teremetskaya, V.P. Varnin, Russ. J. Electrochem. 1995, 31 , 169–174. 46. Y.V. Pleskov, V.Y. Mishuk, M.A. Abaturov, V.V. Elkin, M.D. Krotova, V.P. Varnin, I.G. Teremetskaya, J. Electroanal. Chem. 1995, 396 , 227–232. 47. K. Patel, K. Hashimoto, A. Fujishima, J. Photochem. Photobiol. A: Chem. 1992, 65 , 419–429. 48. K. Patel, K. Hashimoto, A. Fujishima, Denki Kagaku 1992, 60 , 659. 49. R. Tenne, K. Patel, K. Hashimoto, A. Fujishima, J. Electroanal. Chem. 1993, 347 , 409–415. 50. C. Reuben, E. Galun, H. Cohen, R. Tenne, R. Kalish, Y. Muraki, K. Hashimoto, A. Fujishima, J.E. Butler, C. Levy-Clement, J. Electroanal. Chem. 1995, 396 , 233–239.

REFERENCES

15

51. M. Marchywka, P.E. Pehrsson, D.J. Vestyck, D. Moses, Appl. Phys. Lett . 1993, 63 , 3521–3523. 52. R. Ramesham, R.F. Askew, M.F. Rose, B.H. Loo, J. Electrochem. Soc. 1993, 140 , 3018–3020. 53. B. Miller, R. Kalish, L. C. Feldman, A. Katz, N. Moriya, K. Short, A. E. White, J. Electrochem. Soc. 1994, 141 , L41–L43. 54. G. M. Swain, R. Ramesham, Anal. Chem. 1993, 65 , 345–351. 55. G. M. Swain, Adv. Mater. 1994, 6 , 388–392. 56. G. M. Swain, J. Electrochem. Soc. 1994, 141 , 3382–3293. 57. S. Alehashem, F. Chambers, J. W. Strojek, G. M. Swain, R. Ramesham, Anal. Chem. 1995, 67 , 2812–2821. 58. M. Awada, J. W. Strojek, G. M. Swain, J. Electrochem. Soc. 1995, 142 , L42–L45. 59. J.Z. Zhu, S.Z. Yang, P.L. Zhu, X.K. Zhang, G.X. Zhang, C.F. Xu, H.Z. Fan, Fresenius J. of Anal. Chem. 1995, 352 , 389–392. 60. P.L. Zhu, J.Z. Zhu, S.H. Yang, X.K. Zhang, G.X. Zhang, Fresenius J. of Anal. Chem. 1995, 353 , 171–173. 61. J.L. Wu, J.Z. Zhu, L. Shan, N.Y. Cheng, Anal. Chim. Acta 1996, 333 , 125–130. 62. H. B. Martin, A. Argoitia, J. C. Angus, A. B. Anderson, U. Landau, in Applications of Diamond Films and Related Materials: Third International Conference ( A. Feldman et al., eds.), NIST Special Publication 885, U.S. Dept. of Commerce, 1995, pp. 91–94. 63. A. Argoitia, H. B, Martin, E. J. Rozak, U. Landau, J. C. Angus, in “Diamond for Electronic Applications” ( D.L. Dreifus et al., eds.), Mater. Res. Soc. Proc. 1996, 416 , 349–354. 64. H. B. Martin, A. Argoitia, U. Landau, A. B. Anderson, J. C. Angus, J. Electrochem. Soc. 1996, 143 , L133–L136. 65. P. Bouamrane, A. Tadjeddine, J.E. Butler, R. Tenne, C. Levy-Clement, J. Electroanal. Chem. 1996, 405 , 95–99. 66. N. Vinokur, B. Miller, Y. Avyigal, R. Kalish, J. Electrochem. Soc. 1996, 143 , L238–L240. 67. J. Wang, G. M. Swain, M. Mermoux, G. Lucazeau, J. Zak, J. W. Strojek, New Diamond Frontier Carbon Tech. 1999, 9 , 317–343. 68. M.C. Granger, M. Witek, J.S. Xu, J. Wang, M. Hupert, A. Hanks, M.D. Koppang, J.E. Butler, G. Lucazeau, M. Mermoux, J.W. Strojek, G.M. Swain, Anal. Chem. 2000, 72 , 3793–3804. 69. H.B. Martin, A. Argoitia, J.C. Angus, U. Landau, J. Electrochem. Soc. 1999, 146 , 2959–2964. 70. C.H. Goeting, J.S. Foord, F. Marken, R.G. Compton, Diamond Relat. Mat . 1999, 8 , 824–829. 71. C.H. Goeting, F. Marken, A. Gutierrez-Sosa, R.G. Compton, DiamondRelat. Mat . 2000, 9 , 390–396. 72. T.N. Rao, D.A. Tryk, K. Hashimoto, A. Fujishima, J. Electrochem. Soc. 1999, 146 , 680–684. 73. T. Yano, E. Popa, D.A. Tryk, K. Hashimoto, A. Fujishima, J. Electrochem. Soc. 1998, 145 , 1870–1876. 74. I. Yagi, H. Notsu, T. Kondo, D.A. Tryk, A. Fujishima, J. Electroanal. Chem. 1999, 473 , 173–178. 75. T. Yano, E. Popa, D.A. Tryk, K. Hashimoto, A. Fujishima, J. Electrochem. Soc. 1999, 146 , 1081–1087. 76. D.A. Tryk, K. Tsunozaki, T.N. Rao, A. Fujishima, Diamond and Related Mat . 2001, 10 , 1804–1809. 77. A.B. Anderson, D.B. Kang, J. Phys. Chem. A 1998, 102 , 5993–5996.

16

ELECTROCHEMISTRY ON DIAMOND: HISTORY AND CURRENT STATUS

78. I. Yagi, K. Ogai, T. Kondo, A. Fujishima, K. Ueda, K. Uosaki, Chem. Lett . 2003, 32 , 1050–1051. 79. R. Tenne, C. Levy-Clement, Israel J. Chem. 1998, 38 , 57–73. 80. G. M. Swain, A. B. Anderson, J. C. Angus, MRS Bulletin 1998, 23, 56–60. 81. J. C. Angus, H. B. Martin, U. Landau, Y. E. Evstefeeva, B. Miller, N. Vinokur, New Diamond Frontier Carbon Tech. 1999, 9, 175–187. 82. K. Kobashi, ed., Collection of articles from New Diamond and Frontier Carbon Tech. 1999, 9 , Nos. 3 & 5. 83. Y. V. Pleskov, Russ. Chem. Rev . 1999, 68 , 381–392. 84. A. Freedman, C.D. Stinespring, Appl. Phys. Lett . 1990, 57 , 1194–1196. 85. V.S. Smentkowski, J.T. Yates, Science 1996, 271 , 193–195. 86. B. Ohtani, Y.H. Kim, T. Yano, K. Hashimoto, A. Fujishima, K. Uosaki, Chem. Lett . 1998, 9 , 953–954. 87. T.C. Kuo, R.L. McCreery, G.M. Swain, Electrochem. Solid State Lett . 1999, 2 , 288–290. 88. C.H. Goeting, F. Marken, A. Gutierrez-Sosa, R.G. Compton, J.S. Foord, Diamond Relat Mat . 2000, 9 , 390–396. 89. H.J. Mathieu, Surf. Interf. Anal . 2001, 32 , 3–9. 90. G.T. Wang, S.F. Bent, J.N. Russell, J.E. Butler, M.P. D’Evelyn, J. Am. Chem. Soc. 2000, 122 , 744–745. 91. D.R. Fitzgerald, D.J. Doren, J. Am. Chem. Soc. 2000, 122 , 12334–12339. 92. X. Lu, G. Fu, N.Q. Wang, Q.N. Zhang, M.C. Lin, Chem. Phys. Lett . 2001, 343 , 212–218. 93. J.N. Russell, J.E. Butler, G.T. Wang, S.F. Bent, J.S. Hovis, R.J. Hamers, M.P. D’Evelyn, Mat. Chem. and Phys. 2001, 72 , 147–151. 94. J.S. Hovis, R.J. Hamers, M.P. D’Evelyn, J.N. Russell, J.E. Butler, J. Am. Chem. Soc. Comm. 2000, 122 , 732–733. 95. T. Strother, T. Knickerbocker, J.N. Russell, J.E. Butler, L.M. Smith, R.J. Hamers, Langmuir 2002, 18 , 968–971. 96. S. Ida, T. Tsubota, O. Hirabayashi, M. Nagata, Y. Matsumoto, A. Fujishima, Diamond Relat. Mat . 2003, 12 , 601–605. 97. S. Ida, T. Tsubota, S. Tanii, M. Nagata, Y. Matsumoto, Langmuir 2003, 19 , 9693–9698. 98. S.F. Bent, Surface Sci . 2002, 500 , 879–903. 99. J.J. Carey, C.S. Christ, Jr., S.N. Lowery, “Method of Electrolysis Employing a Doped Diamond Anode to Oxidize Solutes in Waste Water,” US patent, 5,399,247, 1995. 100. G. Foti, D. Gandin, C. Comninellis, A. Perret, W. Haenni, Electrochem. Solid State Lett . 1999, 2 , 228–230. 101. M. Fryda, D. Herrmann, L. Shafer, C.P. Klages, A. Perret, W. Haenni, C. Comninellis, D. Gandini, New Diamond Frontier Carbon Tech. 1999, 9 , 229–240. 102. D. Gandini, P.A. Michaud, I. Duo, E. Mahe, W. Haenni, A. Perret, C. Comninellis, New Diamond Frontier Carbon Tech. 1999, 9 , 303–316. 103. D. Gandini, C. Comninellis, A. Perret, W. Haenni, Fifth Eur. Symp. Electrochem. Engng. 1999, 145 , 181–190. 104. P.A. Michaud, E. Mahe, W. Haenni, A. Perret, C. Comninellis, Electrochem. Solid State Lett . 2000, 3 , 77–79. 105. S. Ferro, A. DeBattisti, I. Duo, C. Comninellis, W. Haenni, A. Perret, J. Electrochem. Soc. 2000, 147 , 2614–2619. 106. D. Gandini, E. Mahe, P.A. Michaud, W. Haenni, A. Perret, C. Comninellis, J. Appl. Electrochem. 2000, 30 , 1345–1350.

REFERENCES

17

107. M. Panizza, P.A. Michaud, G. Cerisola, C. Comninellis, J. Electroanal. Chem. 2001, 507 , 206–214. 108. M. Panizza, P.A. Michaud, G. Cerisola, C. Comninellis, Electrochem. Comm. 2001, 3 , 336–339. 109. W. Haenni, J. Gobet, A. Perret, L. Pupunat, P. Rychen, C. Comninellis, B. Correa, Diamond Materials VII, Proceedings (Electrochemical Society Series) 2002, pp. 16–23. 110. J. Iniesta, P.A. Michaud, M. Panizza, G. Cerisola, A. Aldaz, C. Comninellis, Electrochim. Acta 2001, 46 , 3573–3578. 111. M.M. Davila-Jimenez, M.P. Elizalde-Gonzalez, A. Gutierrez-Gonzalez, A.A. Pelaez-Cid, J. Chromatography 2000, 889 , 253–259. 112. C. Terashima, T.N. Rao, B.V. Sarada, D.A. Tryk, A. Fujishima, Anal. Chem. 2002, 74 , 895–902. 113. R. Bellagamba, P.A. Michaud, C. Comninellis, N. Vatistas, Electrochem. Comm. 2002, 4 , 171–176. 114. K. Van Hege, M. Verhaege, W. Verstraete, Electrochem. Comm. 2002, 4 , 296–300. 115. N.S. Lawrence, M. Thompson, C. Prado, L. Jiang, T.G.J. Jones, R.G. Compton, Electroanalysis 2002, 14 , 499–504. 116. Y. Van Andel, L.J.J. Janssen, Appl. Electrochem. 2002, 32 , 481–486. 117. P. Canizares, M. Diaz, J.A. Dominguez, J. Garcia-Gomez, M.A. Rodrigo, Ind. & Eng. Chem. Res. 2002, 41 , 4187–4194. 118. P. Canizares, F. Martinez, M. Diaz, J. Garcia-Gomez, M.A. Rodrigo, J. Electrochem. Soc. 2002, 149 , D118–D124. 119. K.J. McKenzie, D. Asogan, F. Marken, Electrochem. Comm. 2002, 4 , 820–824. 120. E. Brillas, B. Boye, I. Sires, J.A. Garrido, R.M. Rodriguez, C. Arias, P.L. Cabot, C. Comninellis, Electrochim. Acta 2004, 49 , 4487–4496. 121. J.W. Strojek, M.C. Granger, G.M. Swain, T. Dallas, M.W. Holtz, Anal. Chem. 1996, 68 , 2031–2037. 122. Y. Gorbuz, W.P. Kang, J.L. Davidson, D.L. Kinser, D.V. Kerns, Sensors and Actuators B-Chemical 1996, 33 , 100–104. 123. Y. Gorbuz, W.P. Kang, J.L. Davidson, D.V. Kerns, J. Electrochem. Soc. 1999, 146 , 387–391. 124. C.E. Troupe, I.C. Drummond, C. Graham, J. Grice, P. John, J.I.B. Wilson, M.G. Jubber, N.A. Morrison, Diamond Relat. Mat . 1998, 7 , 575–580. 125. T.N. Rao, I. Yagi, T. Miwa, D.A. Tryk, A. Fujishima, Anal. Chem. 1999, 71 , 2506–2511. 126. B.V. Sarada, T.N. Rao, D.A. Tryk, A. Fujishima, New Diamond Frontier Carbon Tech. 1999, 9 , 365–377. 127. B.V. Sarada, T.N. Rao, D.A. Tryk, A. Fujishima, Anal. Chem. 2000, 72 , 1632–1638. 128. E. Popa, D.A. Tryk, A. Fujishima, Diamond Materials VI, Proceedings (Electrochemical Society Series) 2000, 99 , 512–517. 129. A. Fujishima, T.N. Rao, D.A. Tryk, Electrochim. Acta 2000, 45 , 4683–4690. 130. H.B. Martin, P.W. Morrison, Electrochem. Solid State Lett . 2001, 4 , E17–E20. 131. J.K. Zak, J.E. Butler, G.M. Swain, Anal. Chem. 2001, 73 , 908–914. 132. K. Hayashi, Y. Yokota, T. Tachibana, K. Kobashi, T. Fukunaga, T. Takada, J. Electrochem. Soc. 2001, 148 , H17–H20. 133. H. Kawarada, Y. Araki, T. Sakai, T. Ogawa, H. Umezawa, Phys. Stat. Solid. A-Appl. Res. 2001, 185 , 79–83. 134. E. Kohn, M. Adamschik, P. Schmid, A. Denisenko, A. Aleksov, W. Ebert, J. Phys. D-Appl. Phys. 2001, 34 , R77–R85.

18

ELECTROCHEMISTRY ON DIAMOND: HISTORY AND CURRENT STATUS

135. E. Kohn, W. Ebert, M. Adamschik, P. Schmid, A. Denisenko, New Diamond Frontier Carbon Tech. 2001, 11 , 81–100. 136. A.J. Saterlay, J.S. Foord, R.G. Compton, Electroanalysis 2001, 13 , 1065–1070. 137. A. Fujishima, T.N. Rao, Diamond Related Mat . 2001, 10 , 1799–1803. 138. R.L. McCreery, Chem. Rev . 2008, 108 , 2646–2687. 139. O. Chailapakul, W. Siangproh, D.A. Tryk, Sens. Lett . 2006, 4 , 99–109. 140. Y.L. Zhou, J.F. Zhi, Talanta 2009, 79 , 1189–1196. 141. C. Prado, G.U. Flechsig, P. Grundler, J.S. Foord, F. Marken, R.G. Compton, Analyst 2002, 127 , 329–332. 142. T.A. Ivandini, B.V. Sarada, T.N. Rao, A. Fujishima, Analyst 2003, 128 , 924–929. 143. M. Tachiki, Y. Kaibara, Y. Sumikawa, M. Shigeno, T. Banno, K.S. Song, H. Umezawa, H. Kawarada, Phys. Status Solidi A-Appl. Res. 2003, 199 , 39–43. 144. M. Adamschik, M. Hinz, C. Maier, P. Schmid, H. Seliger, E.P. Hofer, E. Kohn, Diamond Relat. Mat . 2001, 10 , 722–730. 145. K. Ushizawa, Y. Sato, T. Mitsumori, T. Machinami, T. Ueda, T. Ando, Chem. Phys. Lett . 2002, 351 , 105–108. 146. W.S. Yang, O. Auciello, J.E. Butler, W. Cai, J.A. Carlisle, J. Gerbi, D.M. Gruen, T. Knickerbocker, T.L. Lasseter, J.N. Russell, L.M. Smith, R.J. Hamers, Nature Materials 2002, 1 , 253–257. 147. T. Knickerbocker, T. Strother, M.P. Schwartz, J.N. Russell, J.E. Butler, L.M. Smith, R.J. Hamers, Langmuir 2003, 19 , 1938–1942. 148. K. Takahashi, M. Tanga, O. Takai, H. Okamura, Diamond Relat. Mat . 2003, 12 , 572–576. 149. S. Wenmackers, K. Haenen, M. Nesladek, P. Wagner, L. Michiels, M. vandeVen, M. Ameloot, Phys Status Solidi A-Appl. Res. 2003, 199 , 4–48. 150. W. Yang, J.E. Butler, J.N. Russell, R.J. Hamers, Langmuir 2004, 20 , 6778–6784. 151. T.L. Lasseter, W. Cai, R.J. Hamers, Analyst 2004, 129 , 3–8. 152. W. Yang and R.J. Hamers, Appl. Phys. Lett . 2004, 85 , 3626–3628. 153. M. Lu, T. Knickerbocker, W. Cai, W. Yang. R.J. Hamers, L.M. Smith, Biopolymers 2004, 73 , 606–613. 154. T. Lasseter, B.H. Clare, B.M. Nichols, N.L. Abbott, R.J. Hamers, Langmuir 2005, 21 , 6344–6355. 155. T.L. Lasseter, B.H. Clare, N.L. Abbott, R.J. Hamers, J. Am. Chem. Soc. 2004, 126 , 10220–10221. 156. B. Rezek, D. Shin, T. Nakamura, C.E. Nebel, J. Am. Chem. Soc. 2006, 128 , 3884–3885. 157. C.E. Nebel, D. Shin, D. Takeuchi, T. Yamamoto, H. Watanabe, T. Nakamura, Diamond Relat. Mat . 2006, 15 , 4–8. 158. D.C. Shin, N. Tokuda, B. Rezek, C.E. Nebel, Electrochem. Comm. 2006, 8 , 844–850. 159. R. Linares, P. Doering, B. Linares, Stud. Health Technol. Inform. 2009, 149 , 284–296. 160. C.E. Nebel, D.C. Shin, B. Rezek, N. Tokuda, H. Uetsuka, H. Watanabe, J. R. Soc. Interf . 2007, 4 , 439–461. 161. V. Vermeeren, S. Wenmackers, P. Wagner, L. Michiels, Sensors 2009, 9 , 5600–5636. 162. M. I. Landstrass, K. V. Ravi, Appl. Phys. Lett . 1989, 55 , 975–977. 163. R.S. Gi, T. Ishikawa, S. Tanaka, T. Kimura, Y. Akiba, M. Iida, Jpn. J. Appl. Phys. 1997, 36 , 2057–2060. 164. R.S. Gi, K. Tashiro, S. Tanaka, T. Fujisawa, H. Kimura, T. Kurosu, M. Iida, Jpn. J. Appl. Phys. 1999, 38 , 3492–3496.

REFERENCES

19

165. R.S. Gi, T. Mizumasa, Y. Akiba, Y. Hirose, T. Kurosu, M. Iida, Jpn. J. Appl. Phys. 1995, 34 , 5550–5555. 166. F. Maier, M. Reidel, B. Mantel, J. Ristein, L. Ley, Phys. Rev. Lett . 2000, 85 , 3472–3475. 167. J. Ristein, M. Riedel, F. Maier, B.F. Mantel, M. Stammler, L. Ley, J. Phys.: Condensed Mat . 2001, 13 , 8979–8987. 168. J.S. Foord, C.H. Lau, M. Hiramatsu, R.B. Jackman, C.E. Nebel, P. Bergonzo, Diamond Relat. Mat . 2002, 11 , 856–860. 169. V. Chakrapani, S.C. Eaton, A.B. Anderson, M. Tabib-Azar, J.C. Angus, Electrochem. Solid State Lett . 2005, 8 , E4–E8. 170. D. Petrini, K. Larsson, J. Phys. Chem. C 2007, 111 , 13804–13812. 171. V. Chakrapani, J.C. Angus, A.B. Anderson, S.D. Wolter, B.R. Stoner, G. U. Sumanasekera, Science 2007, 318 , 1424–1430. 172. V.I. Shapoval, I.A. Novosyolova, V.V. Malyshev, H.B. Kushkhov, Electrochim. Acta 1995, 40 , 1031–1035. 173. A. Denisenko, A. Aleksov, E. Kohn, Diamond Relat. Mat . 2001, 10 , 667–672. 174. J. Ristein, Science 2006, 313 , 1057–1058. 175. P. Strobel, M. Riedel, J. Ristein, L. Ley, Nature 2004, 430 , 439–441. 176. V. Chakrapani, J.C. Angus, A.B. Anderson, G. Sumanasekera, Diamond Electronics Symposium, Mat. Res. Soc. Symp. Proc. 2007, 956 , paper J15-01. 177. V. Chakrapani, C. Pendyala, K. Kash, A.B. Anderson, M.K. Sunkara, J.C. Angus, J. Am. Chem. Soc. 2008, 130 , 12944–12952. 178. S.J. Sque, R. Jones, P.R. Briddon, Phys. Stat. Sol. (a) 2007, 204 , 3078–3084. 179. W. Chen, D. Qi, X. Gao, A.T.S. Wee, Prog. Surf. Sci . 2009, 84 , 279–321. 180. A. Fujishima, Y. Einaga, T.N. Rao, D.A. Tryk, eds, Diamond Electrochemistry, Elsevier/BKS Inc., Tokyo, 2005.

2 Synthesis of Diamond Films Vadali V. S. S. Srikanth and Xin Jiang

2.1

INTRODUCTION

Diamond is a well-known gemstone and considered as a precious possession. About fourth century BC, diamonds were first mined in India, and in the second century BC, Chinese realized diamond’s first industrial use originating from its ultimate hardness. Based on the optical and electrical properties, and the nature and percentage of impurities present, natural diamonds are classified into four types of monocrystalline (types Ia, Ib, IIa, and IIb) and two types of polycrystalline (carbonados and ballas) diamonds. However, type Ia diamonds constitute 98% of the total natural diamonds. The space group for the diamond lattice is O7h (F4,/d 32/m) with two carbon atoms per Bravais cell. Diamond structure can be visualized as two interpenetrating face-centered cubic (FCC) lattices (see Figure 2.1) displaced along the body diagonal by (1/4, 1/4, 1/4)a, where a is the dimension of the cubic unit cell. Each carbon atom of the lattice has a tetrahedral configuration consisting of sp 3 hybrid atomic orbitals. The {111} crystallographic planes are constituted by 6-atom hexagonal rings that are arranged in such a way that the adjacent atoms are alternatively displaced upward and downward from the plane. The stacking sequence along the direction is ABC ABC ABC . . . . In the ˚ same crystallographic direction, the lattice constant and bond length are 3.56 and 1.54 A, respectively. Diamond has two isomers—namely, graphite and lonsdaleite. In graphite, each carbon atom has sp 2 atomic configuration and therefore has three inplane σ bonds; the remaining valence electron forms π bonds using a pz atomic orbital. Consequently, the trigonally bonded 6-carbon rings are situated in a flat plane, contrary to that in the diamond structure. The stacking sequence of the planes in graphite is ABABAB. . . . The lattice constant in ˚ and the in-plane, nearest neighbor the basal plane between repeating layers is 6.707 A, ˚ spacing is 1.42 A. On the other hand, the structure of lonsdaleite is derived from diamond Synthetic Diamond Films: Preparation, Electrochemistry, Characterization, and Applications, First Edition. Edited by Enric Brillas and Carlos Alberto Mart´ınez-Huitle.  2011 John Wiley & Sons, Inc. Published 2011 by John Wiley & Sons, Inc.

21

22

SYNTHESIS OF DIAMOND FILMS

Figure 2.1

FCC structure of the diamond crystal. (Reprinted with permission from Ref. 23.)

A

B

C

A

B

B

A

A

Figure 2.2 Crystal structure of cubic (left) and hexagonal (right) diamond. The difference in stacking sequence of (111) layer pairs in the two structures has been illustrated. (Reprinted with permission from Ref. 23.)

(see Figure 2.2) in that the positioning of atoms in each plane is the same as that in the cubic diamond. However, the stacking sequence of planes is AB AB AB. . . . In lonsdaleite the atoms thus bond closely; the lattice constants in the a and c directions are 2.52 and ˚ respectively. The distance between adjacent atoms is 1.52 A. ˚ 4.12 A, Apart from being a precious gemstone, the diamond has always attracted a special attention due its unique and unsurpassed properties [1–3]. Consequently, research interests have risen to obtain synthetic diamond. Synthetic diamond was obtained for the first time by employing the high-pressure high-temperature (HPHT) technique [4]. In this

2.2 DIAMOND FILM CVD TECHNIQUES

23

technique, graphite and transition-metal (Fe, Ni, Co, etc.,) catalyst powders are mixed and the mixture is subjected to high pressures (∼80–300 kbar) and high temperatures (∼1900–3000◦ C) such that the catalyst component melts. The reaction is then equilibrated for a certain time; the temperature is then gradually decreased to reduce the carbon solubility in the catalyst, resulting in the precipitation of excess carbon as diamond. The role of catalyst in the HPHT technique is not only to lower the activation energy required for the conversion of graphite into diamond but also to dissolve the carbon atoms in the graphite, thus enabling them to reconstitute as diamond. Due to the availability of only expensive natural diamonds and synthetic diamonds only in the form of abrasive grit (via the HPHT technique), it was understood that most properties of diamond could be realized for technological applications only when it is synthesized as a thin film often chemically bonded to an underlying nondiamond substrate material. For almost five decades, chemical vapor deposition (CVD) has been the most commonly used technique to synthesize diamond films. The CVD of diamond involves growth of a tetrahedrally bonded carbon atom network affected by addition of one carbon atom at a time from a suitable gas phase to an initial template; it is accomplished at much lower pressures when compared to the HPHT technique [5–8]. The CVD of diamond thin films has been an interesting topic of research ever since its accomplishment [9–11] and subsequent modifications [12–23]. A typical diamond film growth process via any CVD method includes three main stages that lead to the incorporation of suitable gas species into the growing diamond solid phase: (1) activation of the gas mixtures; (2) gas-phase reactions; and (3) diffusion of gas species onto the substrate surface. Further chemistry occurs on the reaction surface and finally diamond growth takes place. Diamond films’ microstructure can range from oriented columnar grains (including epitaxial grains) to a nanocrystalline structure, depending on the deposition conditions such as surface pretreatment, temperature, pressure, and gas composition. When gas species responsible for diamond growth reach the reaction surface, adhere to it, and settle quickly into possible equilibrium positions before any structural defects form on the growth front, a single crystalline diamond film can be obtained. On the contrary, when the same atoms do not quickly settle into stable equilibrium positions upon their arrival at the reaction surface, nanocrystalline diamond films can be obtained. At present, diamond thin films are serving in a variety of applications related to machining, field emission, electromechanical systems, electrochemical systems, biomedical, biosensing, and others [24–34]. By making a review of diamond film synthesis, the objective of this chapter is to cover the most important scientific and engineering aspects of diamond CVD. This chapter starts with a review of typical diamond thin film synthesis techniques. In the subsequent sections, various aspects pertaining to diamond film (1) nucleation and growth and (2) epitaxy will be discussed. This chapter concludes with a brief discussion of some future perspectives.

2.2 2.2.1

DIAMOND FILM CVD TECHNIQUES History

In the early stages of research concerning CVD of diamond films, the most exciting feature was the formation of the diamond film itself. Although diamond film could be formed, it had been of little technological or commercial importance due to drawbacks

24

SYNTHESIS OF DIAMOND FILMS

such as extremely low growth rate, lack of substrate selectivity, formation of graphite, and other nondiamond forms of carbon. However, with time these drawbacks have been considerably overcome. Early CVD experiments employed mainly thermal decomposition of carbon containing gases at a pressure less than 1 atm to grow a diamond homoepitaxially on natural diamond crystals maintained at 900◦ C. The growth rate of diamonds in such experiments was a meager 1 nm/h; also, graphite was co-deposited alongside diamond. Later it was shown that atomic hydrogen could preferentially etch graphite during the diamond deposition. Subsequent research works extended the possibilities of diamond CVD by showing that diamond could be grown on nondiamond surfaces as well. In the early 1950s, Eversole and Kenmore [9] developed the cyclic pyrolysis process and demonstrated, for the first time, diamond synthesis under low pressures. In the early 1960s, Angus, Will, and Stanko [10] extended Eversole’s work and synthesized a boron-doped diamond film on diamond grit. Eversole’s work was further extended by Derjaguin et al. [11]. Interestingly, in 1953, Schmellenmeier [35] (of Germany) reported synthesis of crystalline diamond containing hard carbon films from acetylene in an electrical discharge. In 1982, Matsumoto et al. [15] made a major breakthrough in diamond CVD technology. A hot filament was used, and hydrocarbon and hydrogen gas mixtures were passed over it to directly activate the gas phase. The diamond film was then deposited onto a nondiamond substrate located about 1 cm away from the filament. Due to the presence of atomic hydrogen during the deposition, graphite was etched away simultaneously; this rendered the previously used deposition and etching cycles unnecessary, which in turn led to higher diamond growth rates. Since then, various gas phase activation methods such as direct current (DC)-plasma, radio frequency (RF)-plasma, microwave plasma, electron cyclotron resonance (ECR)-microwave plasma, and so on, were developed. Mainly, the role of atomic hydrogen in diamond growth was recognized, and the diamond growth rates approached the rates acceptable by industrial standards. In a further development, Rudder, Posthill, and Markunas [36] suggested that the pyrolysis of fluorocarbons such as tetrafluoromethane (CF4 ) could produce epitaxial diamond growth; this suggestion was based on the understanding that OH, O, O2 , F, and F2 species are better graphite etchants than atomic hydrogen. In this method, CF4 and F2 gas mixture diluted in He was blown onto a diamond substrate maintained at 875◦ C. The deposited diamond film was devoid of any graphite. Although the fluorocarbon pyrolysis process took place nearly at thermodynamic equilibrium, a growth rate of only ∼0.6 µm h−1 was achieved. In another development, hydrogen-deficient carbon-containing noble gas plasma was used to deposit phase-pure nanocrystalline diamond films on Si substrates [37]. On the other hand, other non-CVD methods based on low- [38] and high- [39] energy carbon ion beams have also been developed. Methods based on electron bombardment of the substrate have also been developed that resulted in an increase of diamond growth rates almost up to 20 µm h−1 [40,41]. In large areas, as large as 19 cm2 , low pressure diamond film growth has also been reported [42]. Besides CVD, physical vapor deposition methods were also employed [43,44]. A method wherein diamond growth at low temperatures (as low as 75◦ C) was also reported [45]. At present, various developments in plasma-type CVD processes are not only allowing the growth of polycrystalline diamond films with fewer defects but also the growth of single crystalline CVD diamond films with required and highly consistent properties. A good report on the synthesis of diamond in laboratory explaining several of the above discussed progresses is available [46].

2.2 DIAMOND FILM CVD TECHNIQUES

2.2.2

25

Thermal Decomposition Techniques

2.2.2.1 Hot Filament Chemical Vapor Deposition (HFCVD) The hot-filament CVD method is the earliest [15,16] and the most popular method used for diamond film deposition under low pressures (∼20–40 Torr). In this method, a refractory metal (such as tungsten, tantalum, and rhenium) filament is heated to a temperature above 2000◦ C; a methane (CH4 ) and hydrogen (H2 ) gas mixture, typically about 1:99 in volume, is then passed over the hot filament. Hydrocarbon pyrolysis takes place and diamond is deposited on a substrate (at ∼700–900◦ C) kept at a distance of about 8–10 mm from the hot filament. In the HFCVD process, atomic hydrogen is produced, which etches away graphite much faster than diamond. Thus, the deposition rate of diamond is typically about 1 µm h−1 , which is suitable for industrial purposes. For a schematic diagram of typical HFCVD reactor, please refer to [47, p. 155]. In the HFCVD method, the refractory metal filaments are first carburized prior to the deposition of diamond films. The vapor pressure of metal carbides is less than the respective metal. Therefore, at the working filament temperatures, carburization reduces the metallic impurity incorporation in the diamond films. The HFCVD method possesses the ability to adjust to a wide variety of carbon gas sources such as methane, propane, ethane, and other hydrocarbons. Even oxygen-containing hydrocarbons—including acetone, ethanol, and methanol—can be used. The addition of oxygen-containing species widens the temperature range within which diamond deposition can take place. With the aim of improving the phase purity and growth rate of diamond, some modifications have been incorporated into the typical design of HFCVD. In one such modification, a DCplasma is used in combination with the typical HFCVD setup; a bias voltage is applied to the substrate and filament [41,42,48]. A moderate positive voltage is applied to the substrate and a negative voltage to the filament (or an accessory electrode), resulting in electron bombardment of the substrate, which induces desorption of the surface hydrogen which in turn increases the growth rate (up to about 10 µm h−1 ). This technique is called electron-assisted HFCVD. When the bias voltage is high enough to ignite a stable plasma discharge, the decomposition of H2 and hydrocarbon is greatly enhanced, leading to a remarkable increase in the diamond growth rate (up to about 20 µm h−1 ). When the polarity of the bias is reversed—that is, when the substrate is negatively biased—substrates’ surface undergoes ion bombardment, which results in the enhanced nucleation of diamond on nondiamond substrates. Another modification is to replace the single hot filament with multiple filaments or a filament net for uniform diamond film deposition over large areas. 2.2.2.2 Oxy-Acetylene Torch Method Hirose and Kondo [49] were the first to introduce the combustion flame-assisted diamond CVD. In this method, the smoldering tip of a welding torch is used to oxidize a mixture of acetylene (C2 H2 ) and oxygen (O2 ) gases (ratio 1:1). Diamond film gets deposited on the substrate where the tip of the bright interior region of the flame touches the substrate, which is maintained at a temperature of 800–1050◦ C. Under appropriate experimental conditions, the diamond growth rate in this method can go up to 50–100 µm h−1 . Using this method, homoepitaxial films could be deposited at high rates [50]. The aerosol-doping technique, in combination with the oxyacetylene torch method, was used to dope boron into diamond films [51]. When the torch is used in its transversing mode, large-area diamond coatings could be obtained [52]. The advantages of this method over the conventional CVD methods include simplicity and cost-effectiveness of the equipment, lack of power supply, high growth rate, and the

26

SYNTHESIS OF DIAMOND FILMS

ability to deposit diamond over large areas and on curved substrate surfaces. On the other hand, in this method the deposition is difficult to control and consequently the deposited diamond films are inhomogeneous in both microstructure and composition. Also, the torch produces thermal gradients over the substrates’ surface, thereby causing the substrate to warp or fracture during the process of coating large substrate areas. Murakawa, Takeuchi, and Hirose [52] and Cappelli and Paul [53] have achieved great progress in enhancing the diamond film quality and area by using the oxy-acetylene torch method. 2.2.3

Plasma-Aided Deposition Techniques

2.2.3.1 Microwave Plasma-Enhanced CVD (MWCVD) It was found that the atomic hydrogen concentration under CVD conditions could be increased by using a DC plasma ignited by an electrical discharge [14,54]. Thus, plasma became another pathway to dissociate molecular hydrogen into atomic hydrogen and to simultaneously activate hydrocarbon species, thereby leading to the diamond deposition. Besides DC plasma, two other kinds of plasma with different excitation frequencies are possible. The excitation frequency for microwave plasma CVD is typically 2.45 GHz, whereas for radio frequency (RF) plasma CVD it is 13.56 MHz. A schematic of ASTeX microwave plasma CVD reactor is shown in Figure 2.3. Also, Davis [47, p. 159] provides a typical microwave reactor setup.

Microwave Gas inlet

Water cooling Plasma Substrate

Heating

Pump Bias power supply 0 ± 300 VDC

Figure 2.3 Schematic of a microwave plasma-enhanced CVD reactor manufactured by ASTeX. (Reprinted with permission from Ref. 23.)

2.2 DIAMOND FILM CVD TECHNIQUES

27

The microwave frequency used to ignite the plasma oscillates the electrons in the gas mixtures used to deposit diamond. Consequently, high ionization fractions are generated as electrons collide with gas atoms and molecules. Microwave plasma has hot electrons as well as cold ions and neutrals. A substrate typically about 2–3 cm in diameter is placed on a holder in a tube reactor (quartz or steel) compatible with a waveguide that can guide the microwaves generated by a magnetron. Microwaves enter into the reaction chamber from a proprietary antenna that converts a rectangular microwave signal into a circular one. The microwave proceeds through a quartz window into the reaction chamber filled with gas at a particular pressure. The microwave energy is transmitted to the gas mixtures, thereby igniting a luminous ball-shaped plasma, the size of which can be enlarged by increasing the microwave power. The edge of the luminous plasma will be located typically about 2 cm above the substrate’s surface. In this setup, the substrate can be independently heated. Under suitable conditions of gas pressure, microwave power, gas mixture ratio, gas flow rates of different gases, substrate temperature, and so on, diamond film grows on the substrate. 2.2.3.2 DC Plasma CVD Direct current plasma CVD is another deposition technique wherein DC plasma is used to activate the gas source (typically hydrocarbon and hydrogen) for diamond growth. Suzuki et al . [21] developed the DC-plasma-assisted CVD (DC PACVD). In this method, a glow discharge is generated in the reactor between the substrate and the other electrode when the bias voltage is sufficiently high. The activation of the gas phase is thus achieved via the collisions of high-energy electrons with the neutral gas species, resulting in the dissociation of the gas species and the generation of diamond-forming reactive gas species. In all the DC PACVD setups, substrate is mounted on the anode of the reactor. The DC PACVD possesses the ability to coat large areas, which are limited only by the size of the electrodes and the DC power supply. However, this method has a drawback of depositing diamond at very low rates (1 × 1011 /cm3 ) that is favorable for diamond growth. In fact, Hiraki et al. [60] used ECR-MP-CVD to synthesize diamond in 1990. The growth temperature could be reduced to 500◦ C. Later, Yara et al. [61] and Mantei et al. [62] succeeded in diamond deposition using the ECR-MP-CVD technique. They obtained uniform films at substrate temperatures as low as 300◦ C. However, due to the extremely low pressure of the ECR process (10−4 –10−2 Torr), diamond growth proceeds at a very low rate. Therefore, this method is mainly used for laboratory experiments.

2.3 2.3.1

DIAMOND NUCLEATION AND GROWTH Nucleation

2.3.1.1 Definition and Types Diamond nucleation is the first and foremost step in diamond growth via CVD routes. The nucleation process affects film thickness, grain sizes, homogeneity, morphology, defects, surface roughness, and adhesion pertaining to the deposited diamond films. Nuclei are the smallest diamond units that have to be formed for the subsequent diamond growth to start. Nucleation is broadly classified into two types: homogeneous (gas phase) and heterogeneous (on the substrate). There are only few experimental evidences regarding gas phase nucleation [63] of diamond resulting from hydrocarbon molecules such as adamantane, tetracyclododecane, and hexacyclopentadecane acting as the possible diamond embryos. However, it was later shown

2.3 DIAMOND NUCLEATION AND GROWTH

H2 + CH4

Magnetic coil A

2.45 GHz microwave

Plasma

Sample loadlock Magnetic coil B

29

Windows

Diamond film Sample holder

Turbo pump

Figure 2.4 Ref. 23.)

Experimental setup for ECR plasma CVD of diamond. (Reprinted with permission from

that the suggested hydrocarbons are very unstable at CVD temperatures [64]. Additionally, diamond particulates if at all formed in the gas phase were in very low concentrations (when compared to the surface nucleation densities) and would have almost no effect on nucleation [65]. This revelation also led to the idea that small molecular species (for example, CH3 and C2 H2 ), which are capable of surviving the CVD atmosphere and reaching (to be incorporated into the growing diamond nuclei) the seeded substrate, are the ones more suitable for diamond nucleation and growth. On the other hand, heterogeneous nucleation (i.e., surface nucleation) of CVD diamond follows heteroepitaxial growth mechanism from a diamond-seeded (pretreated) substrate surface. First, gas atoms impinge on the reaction surface from the gas phase and get adsorbed onto the surface. Subsequently, the adatoms may either desorb or diffuse over the reaction surface; they may also diffuse into the substrate or bond to other surface atoms. With the deposition time, the surface concentration of the adatoms increases, resulting in the formation of nanoclusters. A statistical fluctuation in the local adatom concentration results in the growth or decay of the clusters. When the clusters attain a size just greater than a critical size, they become stable. Critical size is defined as the size above which growth will be greater than decay of the clusters during the concentration fluctuation. The time taken to form a stable nucleus is called the incubation period . The stabilized

30

SYNTHESIS OF DIAMOND FILMS

clusters then act as suitable sites for further growth either from the continued migration of single adatoms or from direct impingement of atoms from the gas phase. 2.3.1.2 Methods Scratching substrate surface with an abrasive powder has been the most common and powerful method for achieving diamond nucleation densities enough for the formation of continuous diamond films with uniform grain sizes. Mitsuda et al. [66] demonstrated for the first time that scratching the substrate surface with diamond powder could enhance the diamond nucleation density. Besides diamond powder, other abrasive powders such as SiC [67], cubic-BN [68], Cu or stainless steel [69], ZrB2 [70], Al2 O3 [71], and others, have also been used to scratch the substrate surface to enhance the diamond nucleation. Among all the powders, diamond powder has been the most effective one. The scratching technique can be applied to most of the substrates used for diamond growth. For Si substrates, a nucleation density of 107 –108 cm−2 can be routinely obtained after thoroughly scratching it with diamond powder. In contrast, the nucleation density in the case of nonscratched Si substrates can reach only 104 cm−2 . The nucleation density is proportional to the scratching time; the morphology changes from large isolated crystals for short scratching times to smaller, high-density crystals for long scratching times [71,72]. The grit size of the diamond powder used for scratching also influences the nucleation density; 0.25 µm grit is the most effective one for manual scratching [73], and a 40–50 µm powder is best for scratching in an ultrasonic bath using a grit suspension [74]. During scratching with diamond, cubic-BN, or a-SiC powder, the residual powder or fragments are unavoidably left in the scratched grooves and act as seeds for diamond growth. Crystal structures of cubic-BN and a-SiC are close to that of diamond and thereby diamond grows easily on them. In fact, Iijim, Aikawa, and Baba [75] observed the presence of diamond fragments in the scratched grooves of Si, upon which the diamond growth occurred. Scratching with abrasive powder changes the surface morphology; it creates edges, steps, dislocations, and other surface defects. These defects are considered as chemically active sites, which prefer to adsorb diamondforming gas species due to enhanced bonding at high-energy intersecting surfaces with a high density of unsaturated bonds and low coordination numbers [71,76,77]. In this context, using electron microscopy, Denning and Stevenson [70] experimentally observed diamond nucleation on the ridges that have been lithographically etched on nonscratched Si. Singh et al. [71] also reported enhanced nucleation on Si etched by HF/HNO3 . Si+ implantation on nonscratched Si has also been found to enhance the nucleation density [78]. In this method, neither diamond seeds nor carbon-rich layers existed; only the surface morphology or surface structure was changed. However, other attempts to enhance nucleation by creating etch pits on nonscratched Si using acid, H+ [79], or other reactive gas treatments [69] proved unsuccessful. Similarly, attempts to generate large numbers of defect sites via implantation with surface saturation levels of C+ [80] or Ar+ [79] on the unscratched substrates did not result in nucleation enhancement. An array of holes of 0.1–0.3 µm in diameter created by a focused Ga+ ion beam on a Si substrate resulted only in the deposition of nondiamond carbon in the grooves [81]. The reason for the discrepancy among these experiments is unclear. In order to avoid any such problems of understanding, the substrate surface must be thoroughly characterized (with nucleation point of view) prior to any diamond deposition since nucleation is a surface phenomenon and the most critical one. Another method to enhance diamond nucleation is to coat the substrate surface with graphite [82,83], amorphous carbon [82,84], diamond-like carbon [85–87], C60 , and

2.3 DIAMOND NUCLEATION AND GROWTH

31

mechanical oil [84]. This method and the previously mentioned scratching method cannot, however, result in oriented nucleation or epitaxial growth on nondiamond substrates. Bias-enhanced nucleation (BEN) is the method that not only results in higher nucleation densities, but under suitable experimental conditions, results in oriented or epitaxial nucleation. Yugo et al. [88] used the BEN method for the first time and obtained a high nucleation density (109 –1010 cm−2 ) on a mirror-polished substrate (without scratching). In this method a negative bias voltage is generally applied on the substrate in a typical MWCVD setup. BEN can lead to nucleation densities as high as 1010 –1011 cm−2 on mirror-polished Si substrates [89]. Using this method, Jiang and associates [90,91] and Stoner and Glass [92] realized the first ever heteroepitaxial growth of diamond on silicon and silicon carbide substrates, respectively. Yugo, Kimura, and Kanai [93] and Gerber et al. [94] explained the nucleation mechanism by suggesting a shallow ion implantation model in which the sp3 -bonded carbon clusters, formed by low-energy ion implantation, act as the active nucleation sites. The negative bias causes the positively charged ions in the gas phase (plasma in the case of MWCVD) to accelerate toward and bombard the substrate surface, thereby removing any contamination and facilitating the sp3 -bonded carbon cluster formation on the surface. These events in turn enhance the diamond nucleation. On the other hand, Stoner et al. [95] indicated that the critical process for nucleation should be the formation of a surface carbide layer and change in plasma chemistry, such as increase in the concentration of atomic hydrogen caused by substrate biasing. Jiang and colleagues [96,97] found that the overall temporal evolution of the nucleation density corresponds well with a surface kinetic model involving immobile active nucleation sites, germs, and nuclei. They also suggested that, in addition to surface defects (point defects, steps and sp3 -bonded carbon clusters) serving as the nucleation sites, the enhanced surface diffusion and sticking probability of carbon on silicon due to ion bombardment should be the decisive factors. The enhancement of surface diffusion of carbon species has been clearly observed by investigating the distribution of the first nearest-neighbor distances [96]. Stubhan et al. [98] and Lin and associates [40,41] showed that BEN also works in the HFCVD system. In the HFCVD method, the thermally activated gas-phase species (atomic hydrogen and hydrocarbon radicals) are neutral, and as a result, a negative substrate bias as in MWCVD cannot induce enhanced nucleation. However, when plasma is generated by proper choice of bias, enhancement of diamond nucleation similar to that in MWCVD can be achieved. Using BEN in HFCVD, a nucleation density of 109 –1010 cm−2 on mirror-polished Si was obtained, which is similar to the result obtained in the MWCVD system. There are also methods of obtaining high diamond nucleation densities without any surface scratching or substrate biasing. One method is by employing very low gas pressures; the other method is by Si+ implantation [78] into mirror-polished Si substrate prior to the introduction of methane into the deposition chamber. High-density diamond nucleation (109 –1011 cm−2 ) has been obtained on mirror-polished Si substrates using either HFCVD or ECR microwave plasma CVD setup working at very low pressures (0.1–1 Torr) [99]. Similar results have also been obtained on Ti substrates. Using this method, diamond grain density greater than 1010 cm−2 can be obtained, which is two orders of magnitude higher than the highest density (107 –108 cm−2 ) that can be obtained on scratched substrates under conventional CVD process pressure (10–50 Torr) values. In order to obtain high diamond nucleation densities by CVD methods, the substrate surface must be treated in such a way that (1) surface adsorption sites can be created for the incoming hydrocarbon radicals and (2) the distributed region of the adsorption sites

32

SYNTHESIS OF DIAMOND FILMS

is large enough (larger than the critical size of the nucleus) for continuing the growth of diamond nuclei. Si+ ions implantation into the mirror-polished Si substrate meets these criteria exactly. It changes only the surface structure without affecting the composition of Si. After a treatment with Si+ energy of 25 keV (lower energy would be better) and an implantation dose of 2 × 1017 cm−2 , diamond can easily nucleate and grow on a Si wafer. Si+ implantation-enhanced nucleation is assumed to create nanoscale surface defects on the Si substrate. These defects serve as active sites for the adsorption of hydrocarbon radicals necessary for initial diamond nucleation. A similar effect can also be found in the case of diamond growth on porous silicon [100] because, its surface has a dense nanoscale microstructure. 2.3.2

Growth

In order to grow diamond thin films more efficiently and economically, minimize defects in the films, know the most effective gas precursors, and so on, it is critical to understand the diamond film growth mechanism. Tsuda, Nakajima, and Oikawa [101,102] gave the first account of atomic scale {111} diamond growth mechanism [101,102]; it was proposed that diamond growth involved C+ 3 cations or a positively charged surface. However, this mechanism could not be generalized since, under HFCVD conditions, C3+ cations are scarce and the substrate surface is uncharged. Later, Chu et al. [103] proposed that the methyl radicals are the dominant species for the growth of all {111}, {100}, and {110} diamond facets under HFCVD conditions. However, at high temperatures, CH4 and C2 H2 will decompose into various products and it is difficult to distinguish from which original source the products originate. To this end, Martin and Hill [104] and Harris and Martin [105] showed that methyl or methane is more effective than acetylene for diamond film growth. Harris [106] used a nine-carbon model compound [bicyclononane (BCN)] and proposed a growth mechanism involving only neutral CH3 and hydrogen atoms. Frenklach and colleagues [107–110] proposed that acetylene, which is present in greater quantities than CH3 under HFCVD conditions [111–113], is the primary growth specie for {111} diamond and other low-index surfaces. On the other hand, the primary growth specie for nanocrystalline diamond growth has been found to be C2 dimer [114,115] but not CH3 or C2 H2 . In the following (111) and (100) diamond growth mechanisms, and the roles of hydrogen and oxygen gases during diamond film growth will be discussed in detail. Figure 2.5 shows high-resolution electron energy loss (HREEL) spectra obtained in situ during the growth of homoepitaxial (111) diamond surfaces and highly oriented (111) diamond films on Si (0 0 1) for different growth temperatures [116]. The loss peaks at 365, 155, 110, 310, and 460 meV correspond to CH stretching, CH bending, C–C stretching, and the first and second overtones of CH bending vibrations, respectively. The vibration modes and characteristic frequencies of CHx molecular subgroups [117,118] clearly indicate that the surface is not CH3 terminated, but CH terminated. Thus, diamond growth on the (111) surface proceeds via two-by-two layers mode. It can be noted that the CH stretching vibration is perpendicular to the (111) diamond surface, while the CH bending vibration is parallel to the (111) diamond surface. The electron-energy-loss spectral selection rule [119] does not allow the vibration parallel to the (111) surface and therefore should be absent in the spectra. However, in Figure 2.5, CH bending vibration loss peak at 155 meV is present. In order to understand this contradictory presence, the H on the (111) diamond surface was replaced by its isotope D [116]. By doing so, two types of adsorption sites were found to exist; one on the (111) surface, and the other

Intensity (arb. unit)

2.3 DIAMOND NUCLEATION AND GROWTH

33

(a)

(b)

14

(c) –100

100

500 300 Energy loss (meV)

700 ◦

◦

◦

Figure 2.5 HREEL spectra of (1 1 1) diamond facets grown at (a) 800 C, (b) 900 C, and (c) 1000 C. (Reprinted with permission from Ref. 116.)

on another facet. If it is speculated that the growing (111) diamond surface consists of (111) faces and (110) steps, it is not difficult to reconcile the contradiction. The bending vibration of CH on the (110) surface steps is perpendicular to the (111) surface and therefore the CH bending vibration is active. As Frenklach pointed out, growth on the (111) surface takes place in two stages [109,110]. In the first, kernel forms; in the second, it propagates. In the first stage, the appearance of island and the (110) steps are possible. The existence of the CH bending vibration can thus be understood. The second stage of the Frenklach model, which details the connection of C2 H2 in a two-layers by two-layers manner, is experimentally supported [116]. In addition, it has been shown experimentally that during the kernel formation, CH3 is the responsible diamond growth specie [120]. In summary, growth on the (111) diamond surface has been proposed to be completed in two stages. In the first stage, a surface carbon (active site) is activated by H abstraction, adsorption, and catenation of CH3 , which results in kernel formation. In the second stage, the (111) surface grows along the (011) direction aided by acetylene alone. As for the growth mechanism of (100) diamond surface, Harris’ predictions [106] (BCN model) agreed well with experiments. However, the steric repulsion for the H–H site on BCN is different from that on the (100) diamond surface. On the (100) diamond surface, which is terminated by CH2 radicals, a very strong steric repulsion should exist

34

SYNTHESIS OF DIAMOND FILMS

between the neighboring hydrogen atoms because the intermolecular H–H spacing is only ˚ Such a repulsion will ˚ about 0.77 A—nearly the same as that in the H2 molecule (0.74 A). greatly affect the growth of the {100} diamond surfaces. Similar reasons hold true for some low-index surfaces as well. Hamza, Kubiak, and Sulen [112] obtained a 1 × 1 lowenergy electron diffraction (LEED) pattern on the (100) diamond surface at temperatures between 500 and 750 K [121]. This was attributed to the saturation of the dangling bonds of a surface carbon atom by two hydrogen atoms, and thus surface reconstruction did not occur. This surface was a nominally di-hydrogenated surface. Above 1300 K, the surface structure showed a 2 × 1 reconstruction due to desorption of one H atom from a surface carbon atom. This surface was a nominally mono-hydrogenated surface with an elongated C–C dimer bond [122]. A 2 × 1 reconstruction of the (100) diamond surface grown at 1000◦ C was also observed by both scanning tunneling microscopy (STM) and atomic force microscopy (AFM) [123,124]. Figure 2.6 shows HREEL spectra of the (1 0 0) diamond surface grown at a temperature of 800◦ C [125]. Intensity losses occur at 156, 180, and 372 meV, with three smaller losses at 110, 310, and 530 meV. Compared with the characteristic frequencies of the molecular subgroups CHx [115,116], the spectrum is consistent with that of the CH2 radical, which has its stretching, scissors, wagging, twisting, and rocking vibrations at 370, 179, 157, 150, and 108 meV, respectively. In the HREEL spectra, 372, 180, and 110 meV losses are assigned to CH2 stretching, scissors, and rocking vibrations, respectively; and 156 meV loss is assigned to the overlapping of wagging and twisting vibrations. The 310 meV is the overtone of the loss at 156 meV, and the loss at 530 meV is the combination of CH2 wagging and stretching vibrations. The film deposited at 800◦ C exhibits a good crystallinity. At 1000◦ C, the film shows a “cauliflower”-like morphology, and the Raman spectrum exhibits a broad peak at 1580 cm−1 , which is characteristic of graphite, whereas its HREELS is similar to the one at 800◦ C. However, a prominent peak appears at 140 meV. This peak corresponds to the bending vibration of the mono-hydrogenated dimer. Due to the appearance of the mono-hydrogenated surface, hydrogen atoms that are bonded to the (100) surface become further separated. As a result, it releases the strong steric repulsion between the H atoms. Nevertheless, the CH bond of mono-hydrogenated dimer possesses to some extent π bond character. As the hydrocarbon radicals attach to the CH bond, they also take in some π bond character. Their tendency to form graphite is the reason why we observe a broad peak at 1580 cm−1 in the Raman spectrum. Noteworthy is that if the (100) surface of the as-grown sample is exposed to atomic hydrogen (no hydrocarbon involved), the 140 meV loss peak also appears in the HREEL spectrum. This can be explained by the abstraction of one of the two hydrogen atoms bonded to a surface carbon by gas-phase atomic hydrogen. If the abstraction is strong enough, mono-hydrogenated dimers appear in some local regions where the surface is similar to the growth surface at 1000◦ C. If the amount of atomic hydrogen is small, or if the abstraction proceeds at a moderate temperature (∼800◦ C), or if hydrocarbon is involved in the gas source, the abstraction of H will not be strong. Thus, di-hydrogenated carbon atoms remain in the neighborhood of C with one abstracted H, and mono-hydrogenated dimers cannot be formed. The carbon atom then keeps only one H and one vacant site, which may bond to the hydrocarbon radical. As a result, diamond growth proceeds steadily. The quantity of carbon with one hydrogen is determined by the growth temperature, the amount of atomic hydrogen near the surface, and the concentration of activated hydrocarbon. At growth temperatures around 1000◦ C, the (100) diamond surface consists of mono-hydrogenated H–C–C–H dimers as well. For even higher temperatures, a large amount of hydrogen

2.3 DIAMOND NUCLEATION AND GROWTH

35

Intensity (arb. unit)

(a)

(b)

(c) 14

(d) –100

100

300 500 Energy loss (meV)

700 ◦

Figure 2.6 HREEL spectra of (a) (1 0 0) diamond facets grown at 800 C (b) (1 0 0) facets grown at ◦ 1000 C, (c) the sample of (a) dosed with atomic hydrogen, and (d) dosed with oxygen. (Reprinted with permission from Ref. 116.)

desorbs from the surface and the surface carbons become CC dimers. The growth rate of the CH2 -terminated surface is determined by the amount of CH3 radicals in the gas phase. It has been shown that atomic hydrogen can easily abstract one hydrogen atom from CH3 ; therefore, CH2 radicals become the precursor of (100) diamond growth [126]. The control of the appearance of diamond grain facets becomes significant not only in practical usage but also in the testing of established diamond growth mechanisms. The morphology of CVD diamond films is correlated with the growth parameters. Systematic studies on the relationship between the appearances of the (111) and (100) facets and the growth parameters such as the ratio of CH4 /H2 , O2 content, and distance from the hot filament to the substrate have been reported [127,128]. It is well known that, for a kinetically controlled growth system, the crystal morphology is determined by the appearance of facets that have the slowest growth rate in their normal direction and by the corresponding relative growth rates [129]. Because the (110) surface is an S (stepped) face and encounters no repulsion between adjacent hydrogen atoms, it should have the highest growth rate and appear as diamond facets on the film surface. In fact, the growth rate of the (110) surface via the CVD method is the highest among the low-index surfaces, and either CH3 radical- or acetylene-speciesbased mechanisms can be postulated. Thus, in most cases, the surface of CVD diamond crystals appear with {100} faces and {111} faces. According to the established growth

36

SYNTHESIS OF DIAMOND FILMS

model, the growth rate of the (100) face depends on the concentration of CH2 or CH3, whereas the growth rate of the (1 1 1) facet relies on both CH3 and C2 H2 for kernel formation and subsequent growth. If the CH3 concentration near the substrate is much higher than the C2 H2 concentration, the {100} growth rate will be high, and hence the {100} faces will be absent. The crystal appears to have just {111} diamond faces. In contrast, if the C2 H2 concentration near the growth surface is dominant, the {111} diamond face grows faster [130], which consequently results in the appearance of {100} diamond facets. In most cases, the concentrations of CH3 and C2 H2 are comparable, so both {100} and {111} facets are present. Harris and Weiner [131] and Harris and colleagues [132] have measured the dependence of the CH3 and C2 H2 concentrations on the CH4 /H2 and O2 /H2 ratios, as well as on the spacing between the hot filament and substrate. Based on their experimental results, Sun, Zhang, and Lin [125] explained the regularity of the appearance of the diamond facets. In general, CVD diamond growth involves site activation by a surface hydrogen abstraction reaction, (Equation 2.1), followed by addition (Equation 2.2) of a hydrocarbon radical like CH3 [133,134]. Cd represents a surface radical. The competition between surface activation (Equation 2.1) and H-atom recombination with the surface radical site (Equation 2.3) determines the number of active nucleation sites available for a particular set of experimental parameters [133]: Cd H + H∗ ↔ Cd∗ + H2

(2.1)

Cd∗ + CH∗3 ↔ Cd − CH3

(2.2)

+ H ↔ Cd − H

(2.3)

Cd∗

∗

The stable nano-sized crystals formed during the nucleation stage typically exhibit spherical shapes. With time, nucleation density increases to a certain value on which it terminates or ceases to occur at a measurable rate. The isolated crystallites now grow and develop facets due to the relatively high rate of surface carbon diffusion from the surrounding surface sites. Once the crystals grow large enough to coalesce with each another, they form grain boundaries and then continue growing as a continuous film. The morphology of a growing diamond surface depends on the rates at which different diamond planes grow. Under typical growth conditions, the morphology assessment of a diamond film can be made by the growth parameter (Equation 2.4): α=

√ ν100 3 ν111

(2.4)

where ν100 and ν111 are the normal growth velocities of (100) and (111) diamond planes respectively [129,135]. The grains exhibiting fastest growth in the direction perpendicular to the substrate overshadow other slower-growing grains to form a continuous film with a columnar structure [136]. This mode of film growth is known as the evolutionary selection principle [137]. On the contrary to evolutionary selection principle, Jiang et al. [138] showed an oriented diamond growth dependent on facet reactivity and selectivity. Due to the presence of a very little amount of tetramethylsilane gas during a typical MWCVD of diamond, nanocrystalline β-SiC phase selectivity forms on different growing diamond

2.3 DIAMOND NUCLEATION AND GROWTH

37

faces—namely, (001) and (111)—appearing to the gas phase. At the considered experimental conditions, (001) diamond faces avoid any solid deposition resulting from siliconcontaining gas species. In addition, the defective nature of {111} diamond faces will decrease the nucleation barrier and provide high density of active nucleation sites for nano–β-SiC deposition. Figure 2.7 shows the scanning electron microscope (SEM) crosssectional image of the film. There is a single grain growth (marked by arrows) in [001] with its size increasing toward the surface, indicating that tetramethylsilane (TMS) addi¯ direction that tion allows only few suitable nuclei to grow. Typical edges along [110] are formed due to the intersection of {111} planes can also be clearly observed. To explain theoretically this novel growth mechanism, the energy levels of the frontier molecular orbitals (FMOs) and other nearby orbitals of the reactants, a CH3 , a diamond nanoparticle, and a SiH3 , determined at HF/6-31G** level of theory, see Figure 2.8. The growth mechanism was explained by analyzing the reaction occurrence by judging the energy difference between the FMOs of the reactants diamond and CH3 /SiH3 , using frontier orbital theory. The theoretical study revealed that {111} diamond facets are more reactive with SiH3 than the {100} diamond facets, readily forming SiC phases on the {111} facets and leaving the {100} facet clean and with only diamond growth. The selection of diamond (001) plane growth is therefore due to the distortion and/or termination of non-{001} planes growth. Here, the evolutionary selection theory [129,137] is no longer valid because the growth of all the non-{001} faces is interrupted due to nano–β-SiC deposition, even though their growth rates could be higher than that of (001) face. The formation of [001] texture in the present case is therefore due to an angular selection. The {001} facets growth without any tilt will be the fastest and occupy the top-most position, and as a result the tilted facets have no space to develop due to the geometric limitations. Diamond growth models are developed by considering the effects of C1 hydrocarbon radicals—namely, CH3 , CH2 , CH, and C atoms—on both monoradical and biradical reaction surface sites [139]. At typical CVD diamond deposition conditions, diamond growth takes place mainly via the addition of CH3 radicals to monoradical sites. Growth can also take place via the addition of CH3 radicals to the biradical sites; here, the

Figure 2.7 Backscattered field emission scanning electron microscope (FE-SEM) cross-sectional morphology of a film deposited with a continuous variation of TMS gas from a value of 0.0506 to 0% during the deposition. The brighter spot-like regions represent β –SiC phase. (Reprinted with permission from Ref. 138.)

38

SYNTHESIS OF DIAMOND FILMS

0.4 0.3

FMO eigenvalues of reactants (a.u.)

0.2

LUMO {111}

LUMO

{111}

{111}

0.1

LUMO

{100}

0.0 –0.1 {111}

–0.2

{111} {111}

–0.3 –0.4

HOMO {100}

HOMO

HOMO

–0.5 –0.6 –0.7 CH3

Diamond

SiH3

Figure 2.8 The energy levels of FMOs (with isosurfaces as inset) and other nearby orbitals of the reactants of CH3 , diamond nanoparticle (C54 H56 ), and SiH3 . The two energy levels with equal value next to the HOMO for CH3 or SiH3 are degenerate. (Reprinted with permission from Ref. 138.)

unused surface dangling bonds can be rapidly hydrogenated during the conversion of CH3 adducts into a CH2 surface groups [140]. At some deposition conditions, the biradical mechanism is dominant; in this case, the probability of species like C2 , C2 H, and so on, to be the biradical site adducts is enhanced. Such reactive species then have the opportunity to cross-link on the surface, creating a strongly bonded (maybe even nonetchable) defect. This surface defect could act as either a renucleation point for a new epitaxial layer, or, if it is misaligned with the existing lattice, a new crystallite growing in a different direction to that of the main bulk. This last possibility, often termed renucleation, leads to a decrease in the average crystal size. If renucleation occurs frequently, the crystallite size can drop from mm to µm, and eventually to nm, and the films are referred to single crystalline diamond, microcrystalline diamond , and (Ultra) nanocrystalline diamond , accordingly. In addition to CH3 , addition of the other less abundant but highly reactive C1 species, particularly atomic C, to either type of radical site at high H-atom concentration, can also be a route to growth, since the dangling bonds on the adduct would be readily hydrogenated converting the adduct into CH2 . However, at low H-atom concentration, the dangling bonds on the adduct can cross-link to lattice sites, again leading to renucleation and subsequent smaller crystal sizes. Models based on the surface migration of reactive species could predict even diamond growth rates and average crystallite sizes [141,142]. May et al. [143] have critically reviewed the diamond growth mechanism to simulate diamond growth on the (100) surface and presented a unified model for diamond growth based on Monte Carlo simulations [143]. Surface migration, nucleation processes, and

2.3 DIAMOND NUCLEATION AND GROWTH

39

the effects of gas impurities and gas surface reactions have been carefully incorporated into the model to generate diamond growth on the (100) surface. The growth rate is determined mainly by a balance between the flux of CH3 species to the surface and the desorption/etching rate. Using values for the various rates from the literature, a growth rate for standard diamond growth conditions of ∼1 µm h−1 has been estimated, which is consistent with experimental values. Migration of species along and across dimer rows is essential to obtain step-edge growth and smooth surfaces. Without migration, the surfaces become rough and spiky, and the growth rate drops typically by a factor of ∼2, although the exact amount depends on the choice of conditions, and can be several times this value if, for example, a much faster migration rate is used. The growth rate enhancement caused by migration is roughly proportional to the surface diffusion length. It follows that the growth rate enhancement is due to the increased probability of the migrating species meeting a step-edge due to sampling a larger number of surface sites than stationary species. By using a parameter to represent random surface defect formation, growth morphologies resembling “wedding cake” structures can be produced, which again, are consistent with atomic-scale morphologies seen experimentally, and may, in principle, be scaled up to rationalize the cauliflower or ballas diamond morphologies seen at higher C:H ratios. β-scission has been shown to be a minor process, removing only 1-in-800 of the adsorbing carbons from the lattice. Thus, its importance in maintaining a smooth growing surface and in removing longer-chained polymeric species from the surface may have been previously overestimated. 2.3.3

Role of Hydrogen and Oxygen

The most crucial aspect in a typical diamond CVD process is that the hydrocarbon gas must be diluted in hydrogen to as low as 1% and that the hydrogen must be dissociated into atomic hydrogen, which has several specific roles. During the deposition process, atomic hydrogen etches graphite about 20–30 times faster than diamond, resulting in rapid removal of graphite and other nondiamond phases from the substrate and thereby allowing only clusters with diamond structure to remain and grow [125]. The process stabilizes the diamond surface and maintains the sp3 hybridization configuration [144]. Atomic hydrogen not only converts hydrocarbons into the respective radicals that are the necessary species for diamond formation but it also abstracts hydrogen from the hydrocarbons attached on the surface [145], thereby creating active sites for adsorption of the diamond-forming species. However, excess atomic hydrogen causes unnecessarily strong abstraction. The formation of mono-hydrogenated dimers will increase and the graphite phase will readily appear, which results in the deterioration of diamond film quality. As already mentioned, diamond growth is a combined process of deposition and carbon etching that take place concurrently. The growth of diamond occurs if the deposition rate is larger than the etching rate. During CVD growth of diamond films, both atomic H and H+ ions in the plasma cause etching, and H+ ions etch even faster than atomic H [146,147]. It was found that a [001]-textured top layer can be prepared on polycrystalline or [111]-textured diamond films by the application of a negative substrate bias potential during diamond growth [147–149]. An etching effect of hydrogen ions on the growth of diamond films was observed and confirmed to play a dominant role for the [001]-textured growth. The H+ ion bombardment was performed by applying a negative substrate bias during a microwave plasma CVD process, using only hydrogen as the reactant gas. It was discovered that the etching efficiency of H+ ions on non-[001]- oriented grains is

40

SYNTHESIS OF DIAMOND FILMS

higher than that on grains with their (001) faces parallel to the substrate. Lateral growth of the (001) faces can occur during the bombardment process. As a result, the size of the (001) faces increases after H+ etching while grains oriented in other directions are etched off. This effect provides a way to improve the orientation degree of [001]-oriented diamond films and may be helpful for obtaining very thin [001]-oriented diamond films. As for the role of oxygen, its addition (in the form of CO, O2 , or alcohol) to the reaction gases not only has a beneficial effect on the growth rate and quality of the diamond films but it also allows low-temperature diamond growth [150]. Although hydrocarbon species are somewhat reduced by oxygen addition, oxygen has only a relatively small effect on the mole fractions of radical species such as H and CH3 [151]. Also, OH is formed at concentrations sufficient for the removal of nondiamond carbon at a rate comparable to that of diamond growth. At low temperatures ( 0.995) similar to Hg-GC electrode. However, the BDD electrode provides a detection limit three to five times lower than that the Hg-GC electrode because the BDD electrode has a lower background and a higher signal/noise ratio. The detection limits were 50 ppb Zn2+ , 1.0 ppb Cd2+ , 5.0 ppb Pb2+ , 10 ppb Cu2+ , and 1.0 ppb Ag+ . The results demonstrated that BDD is a viable alternative electrode for metal analysis by anodic stripping voltammetry. In the similar concept, the BDD electrodes were also used to investigate the possibility of detecting trace levels of lead by linear-sweep anodic stripping voltammetry [52]. A low limit of detection (2 nM) was obtained in comparison to those by the other electrodes. At low pH values, the Cu2+ concentrations usually present in drinking water do not affect the measurement. Therefore, these methods are attractive and suitable for the determination of trace levels of lead in drinking and tap water. Furthermore, the determination of Zn2+ , Cd2+ , Pb2+ , and Cu2+ by differential pulse anodic stripping voltammetry at BDD electrodes at very low concentrations was reported [53]. Quantification was possible for the simultaneous determination of Zn2+ , Cd2+ , and Pb2+ despite the fact that the Zn2+ and Cd2+ peaks overlapped. Comparison experiments were performed using a glassy carbon electrode, and the BDD electrode provided a lower baseline current, wider potential range, higher sensitivity, and greater longevity than glassy carbon. In 2008, Seehra, Ranganathan, and Manivannan reported the determination of mercury in solutions by BDD electrodes using differential pulse anodic stripping voltammetry [54]. Gold was added to the solution for enhanced sensitivity. Impurity peaks from Cu2+ and Ag+ were identified by choosing the optimal deposition potential. The modified surfaces have been not only used extensively for the analysis of organic and biochemical compounds but some have also been used for inorganic species. Metal and metallic nanoparticles modified boron-doped diamond electrodes and their application to electrochemical analysis have been widely reported. For example, iridium-modified BDD electrodes were used for the electrochemical detection of As3+ by cyclic voltammetry and flow injection analysis with amperometric detection [55]. This method was also applied to the analysis of spiked arsenic in tap water containing a significant amount of various ion elements. Toghill et al. [56] demonstrated a bismuth nanoparticle modified boron-doped diamond (Bi-BDD) electrode for the simultaneous determination of Pb2+ and Cd2+ by square-wave anodic stripping voltammetry (SWASV). In situ plating was achieved using 0.1 mM Bi(NO3 )3 in 0.1 M HClO4 at pH 1.2. The detection limits obtained from this report were 1.9 µg L−1 and 2.3 µg L−1 for Pb2+ and Cd2+ , respectively. Mixtures of As3+ and As5+ at gold-modified diamond electrodes by stripping voltammetry were also electrochemically detected [57]. Good linear responses were observed for standards of As3+ and As5+ . Linear calibration curves were obtained in the concentration range of

166

ELECTROANALYTICAL APPLICATIONS OF DIAMOND FILMS

100–1000 ppb As5+ in a mixture with 100 ppb As3+ (r 2 = 0.99) and in the concentration ranged of 5–30 ppb As3+ in a mixture with 100 ppb As5+ . Detection limits of 5 ppb and 100 ppb were achieved for As3+ and As5+ in a mixture with reproducibility of 7.5% and 8.4%, respectively. This method was used to measure arsenic contaminants in Yokohama tap water.

7.6

FOOD AND DIETARY CONTAMINANTS

Due to the frequent appearance of contaminant compounds from food sample/dietary/supplement matrices, the development of a highly sensitive detection method is needed. Therefore, the use of BDD electrodes as sensors for the electrochemical analysis of food contaminants has been widely published. In 2006, Juliao et al. [58] reported on the electrochemical behavior of nitrofurazone (NFZ) in a predominantly aqueous medium, in the absence and presence of glutathione (reduced form) (GSH), L-cysteine (Cys), and oxygen (O2 ) using BDD electrodes by voltammetry. A linear response was observed for NFZ in the concentration range of 9.9 × 10−7 M to 1.1 × 10−5 M at pH 8.0 with a sensitivity of 2.2 × 106 M cm−2 and a detection limit of 3.4 × 10−7 M. Codognoto et al. [59] demonstrated the examination of electrochemical property of pesticide carbaryl at a BDD electrode with square-wave voltammetry without any previous steps. A linear response was obtained for carbaryl detection from 2.5 × 10−6 M to 30.0 × 10−6 M in 0.1 M Na2 SO4 at pH 6.0 with a detection limit of 8.2 ± 0.2 µg L−1 . Flavonoids, which exhibit high antioxidant activities, were also analyzed in tea samples using BDD amperometric detection coupled to a flow injection system [60]. The experiments were performed at a fixed potential of 0.42 V versus Ag/AgCl with a flow rate of 2.5 mL min−1 . A Britton-Robinson buffer solution at pH 5.0 was used as the carrier stream. A good linear response was observed for a rutin concentration range of 0.1 × 10−4 M to 2.5 × 10−4 M with a detection limit of 7.7 × 10−6 M. In addition, this method was used to determine rutin in three different green teas (Camellia sinensis) with a relative standard deviation of 2.1% (n = 20). In 2008, Medeiros et al. [61] reported the use of square-wave voltammetry in conjunction with a cathodically pretreated BDD electrode for the determination of sodium cyclamate. A linear response was observed for sodium cyclamate in 0.5 M H2 SO4 solution over a concentration range of 5.0 × 10−5 M to 4.1 × 10−4 M and with a detection limit of 4.8 × 10−6 M. The RSD was smaller than 1.2%. In addition, this method was used to determine sodium cyclamate in several dietary products. The use of BDD thin film electrodes to study the electrochemical properties of chloramphenicol using cyclic voltammetry and flow injection analysis was demonstrated [62]. A linear curve for chloramphenicol over the range of 0.1 mM to 10 mM was obtained by cyclic voltammetry. Chloramphenicol was then analyzed by flow injection analysis at −0.7 V versus Ag/AgCl. A linear range of 0.1 µM to 50 µM and a limit of detection of 0.03 µM (S/N = 3) were obtained (see Figure 7.3). This method was successfully applied to the determination of chloramphenicol in sterile eye drops and milk samples using the standard addition method. The average recoveries of chloramphenicol were 98.0% and 93.9 to 103% in eye drops and spiked milk, respectively. N-nitrosamines (N-nitrosopyrrolidine, N-nitrosopiperidine and N-nitrosodiethylamine) in aqueous solutions were also examined at BDD electrodes using cyclic voltammetry

7.6 FOOD AND DIETARY CONTAMINANTS

167

–7.0 Peak current (µA)

(b)

–5.0 Current (µA)

–1.5

(a)

–6.0

(c)

–4.0

y = –0.0216x – 0.0437 R 2 = 0.9948

–1.0 –0.5 0.0 0

(d)

–3.0

(e) –2.0

10

20

30

40

50

60

Chloramphenicol concentration (µM)

(f) (g) (h) (i) (j) (k) (l) (m) (n)

–1.0 0.0

0

500

1000

1500

2000

2500

Time (s) Figure 7.3 Flow-injection analysis with amperometric detection results for various concentrations of chloramphenicol in 0.1 M phosphate buffer (pH 6) in 1% ethanol (a) 1000, (b) 500, (c) 250, (d) 100, (e) 50, (f) 25, (g) 10, (h) 5, (i) 2.5, (j) 1, (k) 0.5, (l) 0.25, (m) 0.1, and (n) 0.05 µM. The calibration graph is also shown in the inset. (Reprinted with permission from Ref. 62.)

[63]. An irreversible oxidation peak was observed at approximately 1.8 V versus Ag/AgCl for all N-nitrosamines. The maximal electrochemical response was obtained using the following square-wave voltammetry parameters: f = 250 Hz, Esw = 50 mV, and Es = 2 mV with a Britton-Robinson buffer solution as the electrolyte (pH 2.0). The detection and quantification limits determined for total N-nitrosamines were 6.0 × 10−8 M and 2.0 × 10−7 M, respectively. Furthermore, the determination of aspartame in dietary products samples without pretreatment was achieved using a BDD electrode [64]. A single irreversible oxidation peak at a potential of 1.6 V versus Ag/AgCl was obtained. A linear response over the range of 9.9 × 10−6 M to 5.2 × 10−5 M with a detection limit of 2.3 × 10−7 M was observed in Figure 7.4. The relative standard deviation (n = 5) was less than 0.2% for a 1.0 × 10−4 M aspartame solution. The proposed method yielded similar results to those obtained from the HPLC method at a 95% confidence level. In 2008, the simultaneous determination of aspartame and cyclamate in dietary products at a BDD electrode was developed [65]. Square-wave voltammetry gave separated oxidation peak potentials in binary mixtures of approximately 400 mV (see Figure 7.5). The detection limits for aspartame and cyclamate were 3.5 × 10−7 M and 4.5 × 10−6 M, respectively, while the relative standard deviations were 1.3% for aspartame and 1.1% for cyclamate. This electrochemical method is simple and highly selective, and can be applied to the determination of aspartame in dietary products. An electrochemical biosensor using a BDD electrode was developed and used for detecting o-nitrophenol released from o-nitrophenyl-β-d-galactopyranose, a reaction catalyzed by β-galactosidase, a marker of E . coli contamination in food [66]. Cyclic voltammogram of o-nitrophenol in a 50 mM phosphate buffer at pH 7.0 showed a well-defined oxidation peak at 0.93 V versus Ag/AgCl. This BDD sensor can be used directly without any surface modification. The enzyme was effectively induced by isopropyl-β-d-thiogalacto-pyranoside. The results showed a biphasic calibration plot

168

ELECTROANALYTICAL APPLICATIONS OF DIAMOND FILMS

12

20

xxxx

8

10

6

9 8

4 2 0

I (µA)

11

10 I (µA)

25

1 2 3 4 5 6 [Aspartame]/10–5 mol L–1

15

7 6 5 4 3

10

2

1

5

0 1.3

1.4

1.5

1.6

1.7

1.8

1.9

E (V) vs. Ag/AgCl Figure 7.4 Square-wave voltammetric response of the BDD electrode for different aspartame concentrations: (1) 0; (2) 9.9 × 10−6 ; (3) 1.5 × 10−5 ; (4) 2.0 × 10−5 ; (5) 2.4 × 10−5 ; (6) 2.9 × 10−5 ; (7) 3.4 × 10−5 ; (8) 3.8 × 10−5 ; (9) 4.3 × 10−5 ; (10) 4.8 × 10−5 ; (11) 5.2 × 10−5 mol L−1 . Insert: Analytical curve for the oxidation process of aspartame. (Reprinted with permission from Ref. 64.)

with a linear range between 4 × 104 cells mL−1 and 2 × 105 cells mL−1 , and 2 × 105 and 6 × 106 cells mL−1 for the first and second regions, respectively. The detection limit was 4 × 104 cells mL−1 with a total analysis time of less than 1.5 h. This method is a suitable alternative for the highly sensitive detection of multiple pathogens simultaneously and can be adapted for field tests to rapidly detect traces of E . coli and other food-borne pathogens. Last, Andrade et al. [67] demonstrated the use of a multidimensional high-performance liquid chromatography method coupled with amperometric detection using a BDD electrode for the simultaneous determination of sulfamethoxazole and trimethoprim in bovine milk. Results showed good linearity in the concentration from 50 to 800 µg L−1 and from 25 to 400 µg L−1 for sulfamethoxazole and trimethoprim, respectively. The intraand interassay coefficients of variation were less than 10% for both drugs. It was also found that LOD values were 25.0 µg L−1 for sulfamethoxazole and 15.0 µg L−1 for trimethoprim. 7.7

MISCELLANEOUS

Several outstanding properties of BDD electrodes make them very attractive for use in many potential applications, including for pH measurements. The electrooxidation of ethylenediaminetetraacetic acid (EDTA) at a thin-film BDD polycrystalline diamond electrode was studied by cyclic voltammetry and amperometry [68]. Under optimal conditions, the linear responses were obtained for concentration ranges from 1.0 × 10−5 M to

7.7 MISCELLANEOUS

169

(a) 80

12 10

60

Ip (µA)

70

8 6

I (µA)

4

50

2 0

40

0

1

2 3 4 –5 –1 [Aspartame]/ 10 mol L

5

30 20

j

10

a

0 1.2

1.4

1.6

1.8

2.0

2.2

E (V) vs. Ag/AgCl (b) 60 30

50 Ip (µA)

40

l 20

k

I (µA)

10

30

0 0

20

1 2 3 –4 –1 [Cyclamate]/ 10 mol L

4

10 0 1.4

1.6

1.8

2.0

E (V) vs. Ag/AgCl Figure 7.5 (a) SW voltammograms for various concentrations of aspartame at a fixed concentration of cyclamate (3.0 × 10−4 mol L−1 in 0.5 mol L−1 H2 SO4 ). Aspartame concentrations (a–j): 5.0 × 10−6 to 5.0 × 10−5 mol L−1 . (b) SW voltammograms for various concentrations of cyclamate at a fixed concentration of aspartame (1.0 × 10−4 mol L−1 in 0.5 mol L−1 H2 SO4 ). Cyclamate concentrations (k–l): 5.0 × 10−5 to 5.0 × 10−4 mol L−1 . Insets are the corresponding analytical curves for the peak current corresponding to the oxidation process of aspartame or cyclamate. (Reprinted with permission from Ref. 65.)

5 × 10−4 M and with a detection limit of 1 × 10−6 M, showing that the BDD electrodes can be used in the quantitative determination of EDTA with a low background current and detection limit. Mitani and Einaga [69] developed the simplest method using BDD electrodes for the analyses of acid concentrations in acidic solutions. Linear sweep voltammetry using BDD electrodes was used to measure the proton concentration. These methods were also applied in vivo for measurements unable of being conducted by conventional glass electrodes.

170

ELECTROANALYTICAL APPLICATIONS OF DIAMOND FILMS

For surface modification, a method for covalent immobilization onto an amineterminated BDD surface was described in 2007 for the determination of phenolic compounds by Zhou and colleagues [70]. The hydrogen-terminated BDD surface was first functionalized by photochemically linking vinyl groups of allylamine to produce a covalently linked amine-terminated active BDD surface. Then the tyrosinase was immobilized onto the active BDD surface using the carbodiimide coupling reaction. The amperometric response was measured as a function of concentration of phenolic compounds in 0.1 M phosphate buffer solution at pH 6.5. The resulted indicated that the BDD enzyme electrode exhibits a good performance in terms of linear dynamic range, sensitivity, and long-term stability to phenolic derivatives as well as the efficient covalent bonding of the enzyme to the substrate and the electrochemical stability of BDD electrode. The tyrosinase-aminophenyl-modified BDD electrode was developed for the detection of phenolic compounds using amperometry. Moreover, Tyrosinase was covalently immobilized on an aminophenyl-modified BDD (AP–BDD) surface via carbodiimide coupling [71]. The effect of the oxygen level, phenolic compound diffusion, and pH of the solution were studied. The Tyr–AP–BDD electrode showed a linear response from 1 to 200 µM, 1 to 200 µM, and 1 to 250 µM and sensitivities of 232.5, 636.7, and 385.8 mA M−1 cm−2 for phenol, p-cresol, and 4-chlorophenol, respectively. Next, the electrochemistry and electrocatalytic activity of cytochrome C (Cyt C) covalently immobilized on a boron-doped nanocrystalline diamond electrode were studied [72]. The linear response to H2 O2 detection was observed over a concentration range of 1 µM to 450 µM, and the detection limit was 0.7 µM. This method showed excellent electrocatalytic performance in terms of fast response, low detection limit, and high stability toward the reduction of H2 O2 . Afterward, Geng et al. [73] described the SiO2 /Cyt C/SiO2 sandwich on the pretreated BDD electrode. Cyt C was immobilized between the SiO2 gel membranes via electrostatic interactions and by carefully controlling the pH of the solution. The SiO2 interlayer was suggested to play a significant role in the sandwich structure of the SiO2 /Cyt C/SiO2 /BDD electrode. The immobilized Cyt C maintained high stability and good electrochemical performance. This electrode was applied to monitor nitrite, and the oxidation current was proportional to the concentration of nitrite in the range of 1.0 × 10−6 M to 1.0 × 10−3 M with a detection limit of 0.5 µM. Finally, the summarization of analytical applications from the use of BDD electrodes for the determination of wide range of analytes is shown in Table 7.1.

7.8

CONCLUSIONS

Since their introduction into the electroanalytical field in the early twentieth century, boron-doped diamond electrodes have become more and more popular. Their unique properties distinguish them from conventional electrode materials, and allow for many electrochemical processes to become more attractive, simple, or even possible. The BDD electrode can be modified with various species such as metals and functional molecules, including biomolecules such as enzymes and nucleic acids. These contribute to their continued progress in new electroanalytical applications. In the future, we expect to see the broad introduction of diamond electrodes in numerous applications.

171

Cyclic and square-wave voltammetry Differential pulse and square-wave voltammetry Cyclic, differential pulse and square-wave voltammetry Cyclic, linear sweep, differential pulse, square-wave voltammetry square-wave voltammetry

Lidocaine

differential pulse voltammetry

Cyclic voltammetry Cyclic and differential pulse voltammetry and chronoamperometry

Sulfonamides (sulfamethoxazole and trimethoprim)

Procaine hydrochloride L-ascorbic acid, Acetaminophen

Sulfonamides (sulfadiazine and sulfamethoxazole)

Pefloxacin

Fluvastatin

Atorvastatin calcium

Flow injection system with pulsed amperometric detection

Three beta-agonists (salbutamol, terbutaline, and clenbuterol)

Technique

Electroanalytical applications of diamond films.

1. Pharmaceutical compound

Analyte

TABLE 7.1

5–200 µM 0.01–0.1 mM

8.01 × 10−6 − 1.19 × 10−4 M (sulfadiazine), 6.10 × 10−6 − 6.01 × 10−5 M (sulfamethoxazole) —

2 × 10−6 − 2 × 10−4 M

9.65 × 10−7 − 3.86 × 10−5 M 1 × 10−6 − 6 × 10−4 M

0.5–100 µM (salbutamol), 1.0–100 µM (terbutaline), 0.5–50 µM (clenbuterol) —

Linear range

2009

3.65 µg L−1 (sulfamethoxazole) and 3.92 µg L−1 (trimethoprim) 0.5 µM —

2008 2008

2008

2.19 × 10−6 M (sulfadiazine), 1.15 × 10−6 M (sulfamethoxa-zole)

2008

2007

2007

2007

10 µg L−1 2.27 × 10−7 1.31 × 10−7 1.37 × 10−7 1.44 × 10−7 4.65 × 10−7 5.78 × 10−7

2006

Year

—

M (DPV) M (SWV) M (DPV), M (SWV) M (DPV) M (SWV)

Limit of detection

AO-BDDE

Comments

(continued overleaf )

16 17

15

14

13

12

11

10

9

Ref.

172 Cyclic and square-wave voltammetry Square-wave voltammetry

Promethazine hydrochloride

Acetylsalicylic acid (ASA)

Biphenyl amino derivatives (2-aminobiphenyl, 3-aminobiphenyl, and 4-aminobiphenyl)

Dopamine in mouse brain Native and thermally denatured fish DNA 2’-deoxyguanosine Purine and pyrimidine bases Proteins

Tryptophan (Trp) and tyrosine (Tyr) Dopamine

Cyclic voltammetry HPLC amperometric detection Cyclic voltammetry and flow injection analysis Amperometric

Differential pulse voltammetry Differential pulse voltammetry Differential pulse voltammetry Cyclic voltammetry

2. Biomolecules or Biological Compounds

Cyclic voltammetry and chronoamperometry

Technique

L-ascorbic acid, Acetaminophen

Analyte

TABLE 7.1 (Continued) Limit of detection

— 26.3–162.1 nM 10 µg mL−1 (BSA), 100 µg mL−1 (IAP) 2 × 10−7 M (2-AB), 3.2 × 10−7 M (3-AB), 5.1 × 10−7 M (4-AB)

10–100 µg mL−1 (BSA), 200–800 µg mL−1 (IAP) 4 × 10−7 −100 × 10−7 M (2-AB), 2 × 10−7 −100 × 10−7 M (3-AB and 4-AB)

—

— — 0.1–10 µM

50 nM

1 × 10−5 M (Trp), 1 × 10−6 M (Tyr) 1.1 × 10−6 M

0.8 and 0.86 µM (L-ascorbic acid), 0.97 and 1.42 µM (Acetaminophen) 2.66 × 10−8 M (peak 1), 4.61 × 10−8 M (peak 2) 2.0 × 10−6 M

0.5 nM − 100 µM

5 × 10−6 M − 1 × 10−4 M

—

2.50 × 10−6 − 1.05 × 10−4 M

5.96 × 10−7 − 4.76 × 10−6 M

0.01–0.1 mM

Linear range

2009

2008

2009 2007

2007

2007

2007

2006

2009

2008

2008

Year

28

27

25 26

24

23

22

21

20

19

18

Ref.

Comments

173

Cyclic voltammetry

Cyclic voltammetry

Dopamine

Dopamine

Phenol (Ph), hydroquinone (HQ) and 4-nitrophenol (4-NP)

4-chlorophenol Oxalic acid

Formaldehyde

3. Pollutant Compounds Cyclic voltammetry and electrochemical impedance spectroscopy square-wave voltammetry Cyclic voltammetry and flow injection analysis Differential pulse voltammetry

Amperometry

Protein

Guanine and adenine bases

Amperometry

Glucose

8 × 10−7 M

5 × 10−6 − 100 × 10−6 M

2006

2006 2006 2007

9.2 µg L−1 0.5 nM 1.82 × 10−6 M (Ph), 1.67 × 10−6 M (HQ), 1.44 × 10−6 M (4-NP) for DPV

0.7–4.0 × 10−5 M 50 nM−10 µM 5 × 10−5 − 1.4 × 10−3 M (Ph), 5 × 10−5 − 7 × 10−3 M (HQ), 5 × 10−5 − 3 × 10−3 M (4-NP) for DPV

2008

—

25 fmol

2009

2008

2008

2006

10–100 mM

—

50 nM

0.3 ng mL−1

1–1000 ng mL−1

—

2.31 × 10−5 M

6.67 × 10−5 − 2 × 10−3 M

37

35 36

34

33

32

31

30

29

(continued overleaf )

GOD/BDD electrode poly-o-ABAmodified BDD AuNPs/polyelec trolyte/BDD electrode Tyramine/PPA/BDD electrode AO-BDDE

174

Oxalic acid Sulfite Oxygen reduction

Thioacetamide

Chemical oxygen demand (COD) COD

Sodium diethyldithiocarbamate 3-nitrofluoranthene and 3-aminofluoranthene

Benzene Aniline

Analyte

TABLE 7.1 (Continued)

Cyclic voltammetry Amperometry

— 0.05 mg L−1 —

0.84 µM

—

2–175 mg L−1

Cyclic voltammetry and flow injection analysis Chronoamperometry 0.02–0.06 mM (in supporting electrolyte 0.1 M Na2 SO4 ) 0.005–0.06 mM (in supporting electrolyte BR2 pH 2.16) — 0.2–20 mg L−1 —

1 mg L−1

20–9000 mg L−1

Amperometry

2008 2008 2008

2008

2009

2007

2007

3 × 10−8 M (3nitrofluoran-thene), 2 × 10−7 M (3-aminofluo ranthene) 7.5 mg L−1

2 × 10−8 − 1 × 10−6 M (3-nitrofluo ranthene), 2 × 10−7 −1 × 10−5 M (3-aminofluorathnene)

Differential pulse voltammetry

2007

2007 2007

8.2 × 10−6 M (Ph), 1.2 × 10−5 M (HQ), 1.1 × 10−5 M (4-NP) for CV — 1 µM

5 × 10−5 − 2 × 10−3 M (Ph), 5 × 10−5 − 1 × 10−2 M (HQ and 4-NP) for CV — 1–5 µM 0.3 µM

Year

Limit of detection

Linear range

—

Cyclic voltammetry Linear-sweep cathodic stripping voltammetry Cyclic voltammetry and chronoamperometry

Cyclic voltammetry

Technique

46 47 48

45

44

43

42

41

39 40

38

Ref.

Py/AuNPs/BDD electrode

Highest Sensitive

Comments

175

Mixtures of As3+ and As5+

Cadmium(II), lead(II)

Arsenite

Hg

Differential pulse anodic stripping voltammetry Cyclic voltammetry and flow injection analysis Square-wave anodic stripping voltammetry Stripping voltametry

Linear-sweep anodic stripping and anodic stripping voltammetry Differential pulse anodic stripping voltammetry

Pb

Zn, Cd, Pb, and Cu

Anodic stripping voltammetry

Cyclic voltammetry and flow injection analysis with amperometry Amperometry and cyclic voltammetric

Zn, Cd, Pb, Cu and Ag

4. Heavy metal

Dissolved oxygen concentration

Oxalate

100–1000 ppb (As5+ ), 5–30 ppb (As3+ )

—

0.1–100 µM

5–20 ppb (Zn), 1.2–25 ppb (Cd), 3.8–45 ppb (Pb), 3–20 ppb (Cu) —

50–1000 ppb (Zn), 1–1000 ppb (Cd), 5–1000 ppb (Pb), 10–1000 ppb (Cu), 1–1000 ppb (Ag) 2–100 nM

—

0.8–100 mM

2008

1.9 µg L−1 (Pb2+ ), 2.3 µg L−1 (Cd2+ ) 100 ppb (As5+ ), 5 ppb (As3+ )

2008

2008

2008

2007

2006

2006

2009

2008

20 nM

1.6 ppb (Zn), 0.36 ppb (Cd), 1.15 ppb (Pb), 0.9 ppb (Cu) 10 ppt

2 nM

50 ppb (Zn), 1 ppb (Cd), 5 ppb (Pb), 10 ppb (Cu), 1 ppb (Ag)

—

32 nM

57

56

55

54

53

52

51

50

49

(continued overleaf )

BiNPs/BDD electrode Au/BDD electrode

Ir/BDD electrode

Allyltriethylammonium Bromide PtNPs/BDD electrode

176 Flow injection analysis with amperometric detection Square-wave voltammetry Cyclic voltammetry and flow injection analysis with amperometric detection Cyclic and square-wave voltammetry Square-wave voltammetry Square-wave voltammetry

Cyclic voltammetry

HPLC-Amperometric detection

Flavonoid

Aspartame Aspartame, cyclamate

E. Coli

Sulfonamides (sulfamethoxazole and trimethoprim)

N-nitrosamines

Sodium cyclamate Chloramphenicol

Cyclic voltammetry Square-wave voltammetry

Technique

Nitrofurazone Carbaryl

5. Food and dietary contaminants

Analyte

TABLE 7.1 (Continued)

2.3 × 10−7 M 3.5 × 10−7 M (aspartame), 4.5 × 10−6 M (cyclamate) 4 × 104 cells mL−1

9.9 × 10−6 − 5.2 × 10−5 M 5.0 × 10−6 − 5.0 × 10−5 M (aspartame), 5.0 × 10−5 − 5.0 × 10−4 M (cyclamate) 4 × 104 − 2 × 105 cells mL−1 (first regions), 2 × 105 − 6 × 106 cells/mL (second regions) 50–800 µg L−1 (sulfamethoxazole) and 25–400 µg L−1 (trimethoprim)

25 µg L−1 (sulfamethoxazole) and 15 µg L−1 (trimethoprim)

6.0 × 10−8 M

2 × 10−6 − 1.36 × 10−5 M

67

66

2008

2009

64 65

63

61 62

60

58 59

Ref.

2007 2008

2008

2008 2008

2006

7.7 × 10−6 M

4.8 × 10−6 M 0.03 µM

2006 2006

Year

3.4 × 10−7 M 8.2 ± 0.2 µg L−1

Limit of detection

5.0 × 10−5 − 4.1 × 10−4 M 0.1 mM to 10 Mm for CV 0.1 µM to 50 µM for FI-amperometry

9.9 × 10−7 − 1.1 × 10−5 M 2.5 × 10−6 − 30.0 × 10−6 M 0.1–2.5 × 10−4 M

Linear range

Comments

177

Cyclic voltammetry and amperometry Linear sweep voltammetry Cyclic voltammetry

Amperometry

Cyclic voltammetry Cyclic and differential pulse voltammetry

Ethylenediaminetetraacetic acid (EDTA) pH in acidic solutions Phenolic compound

Phenolic compound

H2 O2

Nitrite

6. Miscellaneous

0.7 µM 0.5 µM

1.0 × 10−6 − 1.0 × 10−3 M

— 1.0 µM (phenol), 0.5 µM (p-cresol), 0.8 µM (4-cholorophenol) 0.2 µM (phenol), 0.1 µM (p-cresol and 4-cholorophenol)

— 1–175 µM (phenol), 1–200 µM (p-cresol and 4-cholorophenol) 1–200 µM (phenol and p-cresol) 1–250 µM, (4-chlorophenol) 1–450 µM

1 × 10−6 M

1.0 × 10−5 − 5 × 10−4 M

2008

2008

2006

2009 2006

2008

73

72

71

69 70

68

Cyt c/BDD electrode SiO2 /Cyt c/SiO2 / BDD electrode

Tyrosinase/BDD electrode

Tyrosinase/BDD electrode

178

7.9

ELECTROANALYTICAL APPLICATIONS OF DIAMOND FILMS

ACKNOWLEDGMENTS

The authors would like to thank the Thailand Research Fund and Faculty of Science, Chulalongkorn University.

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.

A. Fujishima, T.N. Rao, D.A. Tryk, Electrochim. Acta 2000, 45 , 4683–4690. T.N. Rao, A. Fujishima, Diam. Rel. Mat . 2000, 9 , 384–389. A. Fujishima, T.N. Rao, Diam. Rel.Mat . 2001, 10 , 1799–1803. R.G. Compton, J.S. Foord, F. Marken, Electroanalysis 2003, 15 , 1349–1363. A. Fujishima, in Diamond Electrochemistry (Y. Einaga, T.N. Rao, D.A. Tryk, eds.), Elsevier, Tokyo, 2005. O. Chailapakul, W. Siangproh, D.A. Tryk, Sens. Lett . 2006, 4 , 99–119. O. Chailapakul, E. Popa, H. Tai, B.V. Sarada, D.A. Tryk, A. Fujishima, Electrochem. Commun. 2000, 2 , 422–426. W. Siangproh, in Recent Advances in Analytical Electrochemistry (D. A. Tryk, O. Chailapakul, eds.), Transworld Research Network, India, 2007, pp. 167–197. C. Karuwan, T. Mantim, P. Chaisuwan, P. Wilairat, K. Grudpan, P. Jittangprasert, Y. Einaga, O. Chailapakul, L. Suntornsuk, O. Anurukvorakun, D. Nacapricha, Sensors 2006, 6 , 1837–1850. R.T.S. Oliveira, G.R. Salazar-Banda, V.S. Ferreira, S.C. Oliveira, L.A. Avaca, Electroanalysis 2007, 19 , 1189–1194. B.D. Topal, B. Uslu, S.A. Ozkan, Comb. Chem. High Throughput Screening 2007, 10 , 571–582. B.D. Topal, S. Tuncel, B. Uslu, S.A. Ozkan, Diam. Rel. Mat . 2007, 16 , 1695–1704. B. Uslu, B.D. Topal, S.A. Ozkan, Talanta 2008, 74 , 1191–1200. C.D. Souza, O.C. Braga, I.C. Vieira, A. Spinelli, Sens. Actuators B 2008, 135 , 66–73. L.S. Andrade, R.C. Rocha-Filho, Q.B. Cass, O. Fatibello-Filho, Electroanalysis 2009, 12 , 1275–1280. M. Wei, Y.L. Zhou, J.F. Zhi, D.G. Fu, Y. Einaga, A. Fujishima, X.M. Wang, Z.Z. Gu, Electroanalysis 2008, 20 , 137–143. C. Radovan, C. Cofan, D. Cinghita, Electroanalysis 2008, 20 , 1346–1353. C. Cofan, C. Radovan, Sensors 2008, 8 , 3952–3969. F.W.P. Ribeiro, A.S. Cardoso, R.R. Portela, J.E.S. Lima, S.A.S. Machado, P. de Lima, D. De Souza, A.N. Correia, Electroanalysis 2008, 20 , 2031–2039. E.R. Sartori, R.A. Medeiros, R.C. Rocha, O. Fatibello, J. Braz. Chem. Soc. 2009, 20 , 360–366. G.H. Zhao, Y. Qi, Y. Tian, Electroanalysis 2006, 18 , 830–834. G.H. Zhao, M.F. Li, M.L. Li, Cent. Eur. J. Chem. 2007, 5 , 1114–1123. A. Suzuki, T.A. Ivandini, K. Yoshimi, A. Fujishima, G. Oyama, T. Nakazato, N. Hattori, S. Kitazawa, Y. Einaga, Anal. Chem. 2007, 79 , 8608–8615. A. Apilux, M. Tabata, O. Chailapakul, Bioelectrochemistry 2007, 70 , 435–439. E. Fortin, E. Vieil, P. Mailley, S. Szunerits, T. Livache, Anal. Chem. 2007, 79 , 3741–3746. T.A. Ivandini, K. Honda, T.N. Rao, A. Fujishima, Y. Einaga, Talanta 2007, 71 , 648–655. M. Chiku, T.A. Ivandini, A. Kamiya, A. Fujishima, Y. Einaga, J. Electroanal. Chem. 2008, 612 , 201–207.

REFERENCES

179

28. K. Peckova, K. Jandova, L. Maixnerova, G.M. Swain, J. Barek, Electroanalysis 2009, 21 , 316–324. 29. W. Jing, Q. Yang, Anal. Bioanal. Chem. 2006, 385 , 1330–1335. 30. A. Preechaworapun, A. Suzuki, A. Fujishima, O. Chailapakul, Y. Einaga, Anal. Chem. 2008, 80 , 2077–2083. 31. M. Wei, L.G. Sun, Z.Y. Xie, J.F. Zhi, A. Fujishima, Y. Einaga, D.G. Fu, X.M. Wang, Z.Z. Gu, Adv. Func. Mater. 2008, 18 , 1608–1608. 32. F.J. Shang, Y.L. Liu, S. Hrapovic, J.D. Glennon, J.H.T. Luong, Analyst 2009, 134 , 519–527. 33. S. Hason, H. Pivonkova, V. Vetterl, M. Fojta, Anal. Chem. 2008, 80 , 2391–2399. 34. C.C. Chang, L.C. Chen, S.J. Liu, H.J. Tien, H.C. Chang, J. Chin. Chem. Soc. 2006, 53 , 839–844. 35. V.A. Pedrosa, S.A.S. Machado, L.A. Avaca, Anal. Lett . 2006, 39 , 1955–1965. 36. T.A. Ivandini, T.N. Rao, A. Fujishima, Y. Einaga, Anal. Chem. 2006, 78 , 3467–3471. 37. G.H. Zhao, Y.T. Tang, M.C. Liu, Y.Z. Lei, X.E. Xiao, Chin. J. Chem. 2007, 25 , 1445–1450. 38. Y.Z. Lei, G.H. Zhao, M.C. Liu, X. Xiao, Y.T. Tang, D.M. Li, Electroanalysis 2007, 19 , 1933–1938. 39. R.T.S. Oliveira, G.R. Salazar-Banda, M.C. Santos, M.L. Calegaro, D.W. Miwa, S.A.S. Machado, L.A. Avaca, Chemosphere 2007, 66 , 2152–2158. 40. T. Spataru, N. Spataru, A. Fujishima, Talanta 2007, 73 , 404–406. 41. C. Radovan, F. Manea, Electroanalysis 2007, 19 , 91–95. 42. K. Cizek, J. Barek, J. Fischer, K. Peckova, J. Zima, Electroanalysis 2007, 19 , 1295–1299. 43. H.B. Yu, H. Wang, X. Quan, S. Chen, Y.B. Zhang, Electrochem. Commun. 2007, 9 , 2280–2285. 44. H.B. Yu, C.J. Ma, X. Quan, S. Chen, H.M. Zhao, Environ. Sci. Technol. 2009, 43 , 1935–1939. 45. D. Cinghita, C. Radovan, D. Dascalu, Sensors 2008, 8 , 4560–4581. 46. C.A. Martinez-Huitle, S. Ferro, S. Reyna, M. Cerro-Lopez, A. De Battisti, M.A. Quiroz, J. Braz. Chem. Soc. 2008, 19 , 150–156. 47. C. Chinvongamorn, K. Pinwattana, N. Praphairaksit, T. Imato, O. Chailapakul, Sensors 2008, 8 , 1846–1857. 48. B.E. Roustom, G. Sine, G. Foti, C. Comninellis, J. Appl. Electrochem. 2007, 37 , 1227–1236. 49. T. Kondo, Y. Niwano, A. Tamura, T.A. Ivandini, Y. Einaga, D.A. Tryk, A. Fujishima, T. Kawai, Electroanalysis 2008, 20 , 1556–1564. 50. L. Hutton, M.E. Newton, P.R. Unwin, J.V. Macpherson, Anal. Chem. 2009, 81 , 1023–1032. 51. E.A. McGaw, G.M. Swain, Anal. chim. Acta 2006, 575 , 180–189. 52. D. Dragoe, N. Spataru, R. Kawasaki, A. Manivannan, T. Spataru, D.A. Tryk, A. Fujishima, Electrochem. Acta 2006, 51 , 2437–2441. 53. O. El Tall, N. Jaffrezic-Renault, M. Sigaud, O. Vittori, Electroanalysis 2007, 19 , 1152–1159. 54. M.S. Seehra, S. Ranganathan, A. Manivannan, Anal. Lett . 2008, 41 , 2162–2170. 55. T.A. Ivandini, R. Sato, Y. Makide, A. Fujishima, Y. Einaga, Anal. Chem. 2006, 78 , 6291–6298. 56. K.E. Toghill, G.G. Wildgoose, A. Moshar, C. Mulcahy, R.G. Compton, Electroanalysis 2008, 20 , 1731–1737. 57. D. Yamada, T.A. Ivandini, M. Komatsu, A. Fujishima, Y. Einaga, J. Electroanal. Chem. 2008, 615 , 145–153. 58. M.S.D. Juliao, E.I. Ferreira, N.G. Ferreira, S.H.P. Serrano, Environ. Sci. Technol. 2006, 41 , 5080–5086.

180

ELECTROANALYTICAL APPLICATIONS OF DIAMOND FILMS

59. L. Codognoto, S.T. Tanimoto, V.A. Pedrosa, H.B. Suffredini, S.A.S. Machado, L.A. Avaca, Electroanalysis 2006, 18 , 253–258. 60. V.A. Pedrosa, A.R. Malagutti, L.H. Mazo, L.A. Avaca, Anal. Lett . 2006, 39 , 2737–2748. 61. R.A. Medeiros, A.E. de Carvalho, R.C. Rocha, O. Fatibello, Quim. Nova 2008, 31 , 1405–1409. 62. S. Chuanuwatanakul, O. Chailapakul, S. Motomizu, Anal. Sci . 2008, 24 , 493–498. 63. R.T.S. de Oliveira, G.R. Salazar-Banda, S.A.S. Machado, L.A. Avaca, Electroanalysis 2008, 20 , 396–401. 64. R.A. Medeiros, A.E. de Carvalho, R.C. Rocha, O. Fatibello, Anal. Lett . 2007, 40 , 3195–3207. 65. R.A. Medeiros, A.E. de Carvalho, R.C. Rocha-Filho, O. Fatibello-Filho, Talanta 2008, 76 , 685–689. 66. E. Majid, K.B. Male, J.H.T. Luong, J. Agric. Food. Chem. 2008, 56 , 7691–7695. 67. L.S. Andrade, M.C. de Moraes, R.C. Rocha-Filho, O. Fatibello-Filho, Q.B. Cass, Anal. chim. Acta. 2009, 654 , 127–132. 68. A.E. Denisova, Y.V. Pleskov, Russ. J. Electrochem. 2008, 44 , 1083–1085. 69. N. Mitani, Y. Einaga, J. Electroanal. Chem. 2009, 626 , 156–160. 70. Y.L. Zhou, R.H. Tian, J.F. Zhi, Biosens. Bioelect . 2007, 22 , 822–828. 71. Y.L. Zhou, J.F. Zhi, Electrochem. Commun. 2006, 8 , 1811–1816. 72. Y.N. Zhou, J.F. Zhi, Y.S. Zou, W.J. Zhang, S.T. Lee, Anal. Chem. 2008, 80 , 4141–4146. 73. R. Geng, G.H. Zhao, M.C. Liu, M.F. Li, Biomaterials 2008, 29 , 2794–2801.

8 Cathodic Pretreatment of Boron-Doped Diamond Electrodes and Their Use in Electroanalysis Leonardo S. Andrade, Giancarlo R. Salazar-Banda, Romeu C. Rocha-Filho, and Orlando Fatibello-Filho

8.1

INTRODUCTION

Diamond electrodes have been a subject of investigation and application ever since 1987, when Pleskov et al. [1] reported pioneering electrochemical measurements carried out on undoped diamond films. Most of the early work on diamond electrodes, both fundamental and applied, was summarized in the book edited by Fujishima et al. [2], in book chapters by Angus, Pleskov, and Eaton [3] and by Swain [4,5], and in short reviews on boron-doped diamond (BDD) sensors by Chailapakul, Siangproh, and Tryk [6], Chen [7], and Park et al. [8]. More recently, Peckov´a, Musilov´a, and Barek [9] reviewed the voltammetric determination of organic substances using BDD electrodes, whereas Luong, Male, and Glennon [10] reviewed their functionalization and analytical applications. The electrochemical behavior of BDD electrodes depends on their physical, chemical, and electronic properties, which can be significantly affected by the surface termination (hydrogen, oxygen, and others). Microcrystalline BDD films currently prepared by

Synthetic Diamond Films: Preparation, Electrochemistry, Characterization, and Applications, First Edition. Edited by Enric Brillas and Carlos Alberto Mart´ınez-Huitle. © 2011 John Wiley & Sons, Inc. Published 2011 by John Wiley & Sons, Inc.

181

182

CATHODIC TREATMENT OF BDD ELECTRODES AND THEIR USE IN ELECTROANALYSIS

chemical vapor deposition (CVD) methods are hydrogen terminated (HT-BDD). Many redox couples present relatively high electron transfer rates on HT-BDD. In this sense, in 1997, in one of the first efforts to use BDD thin-film electrodes in electroanalysis, Jolley et al. [11] reported reversible to quasi-reversible electron transfer kinetics for the hexacyanoferrate(II)/hexacyanoferrate(III) [Fe(CN)6 4−/3− ] redox couple on a CVD-prepared BDD electrode that did not undergo any surface pretreatment before use. Afterward, Granger and Swain [12] reported on the influence of oxygen and hydrogen surface terminations of BDD electrodes (as deposited and obtained by plasma treatments) on the electron transfer involving Fe(CN)6 4−/3− and other redox couples. Since the electron transfer was much faster on HT-BDD electrodes, the researchers concluded that the Fe(CN)6 4−/3− redox reaction proceeds through a specific surface interaction available on the hydrogen-terminated surface and that such surface interactions appear to be blocked on the oxygen-terminated surface. However, they observed only a quasireversible response (Ep = 70 mV) for the hydrogen-terminated surface. The surface termination on BDD surfaces is usually generated by electrochemical methods (hydrogen evolution to produce hydrogen terminations; oxygen evolution for oxygen terminations) or RF-plasma treatment (H and O terminations) among others [13–28]. Hydrogen-terminated surfaces are hydrophobic with a negative electron affinity and are highly conductive [29–32], whereas oxygen-terminated ones are hydrophilic, with a positive electron affinity, and present a low conductivity [18,33,34]. As a consequence, the charge transfer rate for some redox couples, including Fe(CN)6 4−/3− , can significantly change with the surface termination on the BDD electrode. Thus, many authors have investigated the charge transfer mechanism and kinetics of several redox couples on BDD electrodes, using, for this purpose, physical and electrochemical methods [12,18, 35–45]. Here, we will concentrate on the cathodic pretreatment of BDD electrodes, which produces HT-BDD electrodes. First, investigations on the effect of the cathodic pretreatment of BDD electrodes on their electrochemical properties will be presented. Second, the use of HT-BDD electrodes on the determination of specific analytes in different matrixes will be reviewed. Third, the effect of the electrochemical pretreatments on the deposition of metals on BDD electrodes will be briefly reviewed.

8.2 CATHODIC PRETREATMENT OF CONDUCTIVE DIAMOND FILMS Despite the fact that enhanced electron transfer kinetics for some redox couples and high conductivity for HT-BDD surfaces (as-prepared or obtained by plasma treatments) were reported since the latter 1990s [12], in 2004, Suffredini et al. [18] called attention to the enhanced electrochemical response of BDD electrodes brought on by a cathodic surface pretreatment, for a variety of reactions. This research was inspired by previous observations in the analytical detection of pentachlorophenol (PCP) and 4-chlorophenol (4-CP) using square-wave voltammetry (SWV) at cathodically [−3.0 V versus Ag/AgCl (3.0 mol L−1 KCl) for 30 s] pretreated BDD electrodes [43,45], when excellent analytical performances were obtained. Thus, Suffredini et al. [18] compared the effect of anodic and cathodic electrochemical pretreatments [applying ±3.0 V versus Ag/AgCl (3.0 mol L−1 KCl), for 30 min, in 0.5 mol L−1 H2 SO4 ] on the electrochemical response

8.2 CATHODIC PRETREATMENT OF CONDUCTIVE DIAMOND FILMS

183

(a) 20

I (µA)

15

10

5

0 0.5

0.6

0.7 0.8 E (V) vs Ag/AgCl

0.9

1.0

8 (b)

I (µA)

6

4

2

0

0.2

0.4

0.6 0.8 E (V) vs Ag/AgCl

1.0

1.2

Figure 8.1 Cyclic voltammetry (ν = 0.05 V s−1 ) on BDD (A = 0.62 cm2 ) for 5.0 × 10−5 mol L−1 pentachlorophenol (a) and 4-chlorophenol (b) in a 0.1 mol L−1 Britton-Robson buffer (pH 5.5), after anodic (dotted lines) and cathodic pretreatments (full lines). (Reprinted with permission from Ref. 18.)

for the oxidation of these analytes. As seen in Figure 8.1, in both cases the electroanalytical response was significantly enhanced at the cathodically pretreated BDD electrode. Suffredini et al. [18] also reported on the effect of the electrochemical pretreatments of BDD surfaces on the cyclic voltammetry and electrochemical impedance spectroscopy (EIS) responses for solutions containing K4 [Fe(CN)6 ] in aqueous 0.5 mol L−1 H2 SO4 . The results of the cyclic voltammetry experiments show that the hexacyanoferrate(II)/hexacyanoferrate(III) redox couple behaves in a quasi-reversible manner after the anodic pretreatment, whereas the cathodic pretreatment leads to a reversible behavior with a Ep value of 60 mV. In addition, the EIS data for this redox couple showed that the value of a resistance associated to a high-frequency element presented values of over 300  cm2 and about 4  cm2 after anodic and cathodic pretreatments, respectively (see Figure 8.2). From these studies the authors concluded that the electrochemical response of BDD electrodes is strongly affected by the type of electrochemical pretreatment applied to their

184

CATHODIC TREATMENT OF BDD ELECTRODES AND THEIR USE IN ELECTROANALYSIS

(a) 400

−Z ′′ (Ω cm2)

320

240

160 0.15 Hz 80 5.0 Hz 0 0

80

160

240 Z ′ (Ω

400

(b)

12

−Z ′′ (Ω cm2)

320

cm2)

8

4 320 Hz

0 0

4

8 Z ′ (Ω

cm2

12

)

Figure 8.2 Complex-plane plots for the BDD electrode in a 1.0 mmol L−1 K4 [Fe(CN)6 ] + 0.5 mol L−1 H2 SO4 aqueous solution after anodic (a) or cathodic (b) pretreatment. Frequency intervals: 50 kHz to 30 mHz (a) and 50 kHz to 10 Hz (b). Measurements carried out at 0.60 V versus HESS—hydrogen electrode in the same solution. (Reprinted with permission from Ref. 18.)

surface before measurements. This effect is very noticeable for some analytes, although it was also present for electron transfer reactions of well-known reversible couples, for which a cathodic pretreatment of the BDD electrode surface prior to measurements leads to an enhanced electrochemical activity. After, Mah´e, Devilliers, and Comminellis [19] also reported data that evidenced a significantly improved electrochemical reactivity of the Fe(CN)6 4−/3− redox couple on

8.2 CATHODIC PRETREATMENT OF CONDUCTIVE DIAMOND FILMS

185

cathodically pretreated BDD electrodes. According to these authors, the cathodic polarization gives rise to a very high rate constant with great reproducibility, which is consistent with the existence of a hydrogen participation into the hole-generation process at BDD surfaces. They also concluded that simple electrochemical treatments could be used to achieve reversible chemical surface modification of BDD electrodes, associated to a drastic change in the rate of electron transfer. In 2006, Salazar-Banda et al. [20] reported on studies to elucidate the effect of different pretreatments (anodic, cathodic, and thermal) on 800 ppm BDD electrodes on the electrochemical response of the hexacyanoferrate(II)/hexacyanoferrate(III) redox couple. As-prepared BDD electrodes were subjected to either an anodic (3.0 V versus HESS, for 30 min) or a cathodic (−3.0 V versus HESS, for 3 or 30 min) pretreatment in a 0.5 mol L−1 H2 SO4 aqueous solution; the thermal pretreatment was carried out at 400◦ C during 30 min in an oven with an oxygen atmosphere. The kinetics of the Fe(CN)6 4−/3− redox couple was greatly affected by the different pretreatments, as can be seen in Figure 8.3. Clearly, the cathodic pretreatment significantly facilitated the redox reaction, leading to a reversible behavior with a Ep of 60 mV and an Ipox /Ipred ratio of 1.0. An analysis of baseline for the cathodically pretreated electrode, also shown in Figure 8.3, concludes that HT-BDD presents a clean surface with no faradaic processes except for a small anodic signal at the positive end of the scan; this signal has been attributed to the presence of small amounts of sp2 carbon as surface impurity [19,46]. On the other hand, the anodic and thermal pretreatments yielded electrode surfaces that do not favor the kinetics of this redox reaction, resulting in irreversible behaviors, with Ep values of 1010 mV and 440 mV associated to Ipox /Ipred ratios equal to 1.5 and 1.3, respectively. This clearly confirms that the cathodic pretreatment of the BDD electrode (resulting in a hydrogen-terminated surface) leads to an enhanced electrochemical response, as previously reported [18].

Figure 8.3 Cyclic voltammograms (ν = 0.05 V s−1 ) for BDD electrodes (A = 0.63 cm2 ) in the presence of a 1.0 mmol L−1 K4 [Fe(CN)6 ] + 0.5 mol L−1 H2 SO4 aqueous solution after cathodic, anodic, and thermal pretreatments. Also included is the background response for the cathodically pretreated electrode. (Reprinted with permission from Ref. 20.)

186

CATHODIC TREATMENT OF BDD ELECTRODES AND THEIR USE IN ELECTROANALYSIS

Salazar-Banda et al. [20] also investigated the effect of time of cathodic pretreatment on the electrochemical response of the pretreated BDD electrode. They showed that the voltammetric behavior of the Fe(CN)6 4−/3− redox reaction is highly dependent on the duration of the cathodic pretreatment. As seen in Figure 8.4, the reversibility of this redox reaction increases significantly even after pretreatment times as short as 3 s. The value of Ep decreases from 420 mV for the as-received electrode to 110 mV, 72 mV, and 60 mV after pretreatment times of 3 s, 3 min, and 30 min, respectively. Thus, a pretreatment time of 30 min (corresponding to a passed charge of about −60 C cm−2 ) led to a totally reversible electrochemical response of the Fe(CN)6 4−/3− redox couple. It is worth highlighting that short pretreatments time and, consequently, low charges passed already had a significant effect on the BDD electrochemical response. Suffredini et al. [18] attributed the large differences in the electrochemical response of BDD electrodes after cathodic and anodic pretreatments to either an internal transformation of the BDD film or, most probably, to the presence of a discontinuous passive layer. However, Raman studies performed on 800 ppm BDD electrodes before and after a cathodic pretreatment showed that neither important bulk structural differences nor significant changes in the sp2 /sp3 carbon content are introduced into the BDD film by the pretreatment [20]. Thus, this indicates that the enhanced electrochemical response brought on by the cathodic pretreatment is due only to superficial changes, most probably in the surface terminations of the BDD film that becomes mainly hydrogen terminated. In this sense, it is important to mention that several studies have reported that the surface of nondoped diamond is either conducting when it is hydrogen terminated or insulating when oxygen terminated [31,47–51]. Furthermore, the hydrogen-terminated surface becomes insulating when heated under vacuum at temperatures higher than 700◦ C [52], a process that leads to the elimination of hydrogen from the diamond surface.

tcp = 3min 60

tcp = 30min

I (µA)

30 tcp = 0 0

–30 tcp = 3 s –60 0.2

0.4

0.6 E (V) vs. HESS

0.8

1.0

Figure 8.4 Effect of cathodic pretreatment time on the electrochemistry of BDD electrodes (A = 0.63 cm2 ): cyclic voltammograms (ν = 0.15 V s−1 ) for an 800 ppm BDD electrode in a 1.0 mmol L−1 K4 [Fe(CN)6 ] + 0.5 mol L−1 H2 SO4 aqueous solution after cathodic pretreatments at −3.0 V versus HESS, for different lengths of time (tcp ), as indicated on the curves. (Reprinted with permission from Ref. 20.)

8.2 CATHODIC PRETREATMENT OF CONDUCTIVE DIAMOND FILMS

187

However, the conductivity is reinstated when the diamond is hydrogenated again by a plasma treatment. According to Hayashi et al. [29], the conductivity of nondoped hydrogen-terminated diamond is in the range of 10−6 to 10−4 S, due to p-type carriers with a lateral concentration in the range of 1012 to 1013 cm−2 and with a mobility between 10 and 100 cm2 V−1 s−1 . The mobility measured for BDD is not too different, there being a general agreement that the carriers are holes residing in a layer near to the surface [50,51]. Thus, on nondoped diamond, a hydrogen-terminated surface is necessary for the existence of this superficial conductivity that allows the onset of bulk conductivity. On the other hand, if the diamond surface is insulating (oxygen terminated), no bulk conductivity arises. Taking this into account, in electrodes with low boron doping levels it could be assumed that there is a mixed conductivity at the BDD electrode surface, due to areas with high boron contents and to hydrogen-terminated areas. These areas must be connected to boron-rich regions in the diamond bulk, leading to conductive pathways [35]. Probably one of the most interesting results reported by Salazar-Banda et al. [20] is the one that cathodically pretreated BDD electrode surfaces present a dynamic electrochemical behavior; that is, the HT-BDD electrochemical response changes with time of exposition to air. Hence, after a cathodic pretreatment (−3.0 V versus HESS, for 30 min), the HT-BDD electrodes present a progressive decrease of the electron transfer rate for the Fe(CN)6 4−/3− redox couple that results in a loss of the reversibility as a function of time exposed to atmospheric conditions. This dynamic behavior was associated to a loss of superficial hydrogen due to oxidation of the surface by oxygen from the air. This assumption was confirmed by XPS analysis, since a cathodically pretreated electrode exposed to air for 30 days clearly presented an increased superficial content of oxygen. For a cathodically pretreated 800 ppm BDD electrode, Ep changes from 60 mV to 85 mV, after the first 48 h of exposition to the atmosphere, and to 293 mV after 100 days (see Figure 8.5) [20]. These changes must be related to loss of superficial hydrogen due to the oxidation of the surface by atmospheric oxygen or by other species (HCO3 − , OH− ) contained in the thin layer of water naturally formed on the surface of solids exposed to air [31,50]. In fact, when these results were correlated with oxygen content from XPS data, it was possible to observe that after one month of exposition to air, an increase of 85 mV in the value of Ep is related to a difference of 5.8% in the value of the O/C ratio on the electrode surface [20]. In this sense, Kulesza, Patyk, and Rozploch [53] reported on the spontaneous oxidation of hydrogenated nondoped diamond surfaces. When hydrogenated nondoped diamond surfaces are in contact with air, a degradation of the hydrogen layer occurs by oxidation, followed by long-term deterioration of the surface electrical conductivity. The increase rate of the surface resistance was calculated to be approximately (10 ± 6) k month−1 . If one assumes that the surface resistance of BDD electrodes exposed to air increases similarly to the one for hydrogenated nondoped diamond surfaces, this may explain the variations in the electrochemical behavior of the hexacyanoferrate(II)/hexacyanoferrate(III) redox couple (see Figure 8.5). Additionally, Salazar-Banda et al. [20] reported that the dynamic behavior of the electrochemical response of cathodically pretreated BDD electrodes also presents an inverse dependence with the doping level. Figure 8.6 shows how the value of Ep changed with time of exposition to air for the different cathodically pretreated (−3.0 V versus HESS, for 30 min) BDD electrodes. From this figure one can see, for instance,

188

CATHODIC TREATMENT OF BDD ELECTRODES AND THEIR USE IN ELECTROANALYSIS

Figure 8.5 Effect of exposure to air on the electrochemical response of a cathodically pretreated 800 ppm BDD electrode (A = 0.63 cm2 ): cyclic voltammograms (ν = 0.05 V s−1 ) for the Fe(CN)6 4−/3− redox couple as a function of the time after pretreatment (tap ) that the BDD electrode was exposed to atmospheric conditions, recorded in a 1.0 mmol L−1 K4 [Fe(CN)6 ] + 0.5 mol L−1 H2 SO4 aqueous solution. Cathodic pretreatment: −3.0 V versus HESS, for 30 min, in 0.5 mol L−1 H2 SO4 . (Reprinted with permission from Ref. 20.)

Figure 8.6 Effect of exposure to air on the Ep values for the Fe(CN)6 4−/3− redox couple for cathodically pretreated (−3.0 V versus HESS, for 30 min) BDD electrodes having different doping levels, as indicated in the inset. Data from cyclic voltammograms (ν = 0.05 V s−1 ) recorded in a 1.0 mmol L−1 K4 [Fe(CN)6 ] + 0.5 mol L−1 H2 SO4 aqueous solution. (Reprinted with permission from Ref. 20.)

8.2 CATHODIC PRETREATMENT OF CONDUCTIVE DIAMOND FILMS

189

that after 140 h of exposition to the atmosphere the Ep values for the 300 ppm, 800 ppm, 2000 ppm, and 8000 ppm BDD electrodes increased by 35, 32, 8, and 5 mV, respectively. This dynamic behavior of the electrochemical response inversely dependent on the electrode doping level clearly suggests that the boron content has a stabilizing effect on the H-terminated surface. According to the authors, these results can be further understood if one assumes that there is a mixed conductivity at the surface of the BDD electrodes with low boron doping levels, due to the existence of heterogeneity in the boron content on the surface; the boron-poor areas lose hydrogen quicker than the boron-rich areas, due to a stabilizing effect of boron. These areas must be connected to boron-rich regions in the diamond bulk, leading to conductive pathways [35]. However, this mixed conductivity is less important in electrodes with high boron doping levels, where most of the surface is covered with boron-rich areas (hydrogen stabilized). The findings of the study by Salazar-Banda et al. [20] clearly show that hydrogenterminated sites play an important role in the electrochemical response of BDD electrodes. Findings also indicate that when reproducible results are required, the BDD electrode has to be cathodically pretreated just before the electrochemical experiments are carried out, especially when the electrode was not used for a long period of time. Girard et al. [21] studied the influence of anodic and cathodic pretreatments on the behavior of highly doped (1020 cm−3 ) BDD electrodes. The electrochemical pretreatments were performed during 10 s with two significantly different current densities: ±10−4 A cm−2 (±10−3 C cm−2 ) for mild pretreatments and ±10−1 A cm−2 (±1 C cm−2 ) for severe ones. The authors observed that when the mild cathodic pretreatment (−10−4 A cm−2 ) was carried out, the reversibility of the Ce3+/4+ and Fe(CN)6 4−/3− redox couples diminished (see Figure 8.7). It is worth noticing that the mild cathodic pretreatment is indeed very mild, associated to a potential difference between working and reference electrodes that is nearly stable and about −1.0 V versus MSE [mercury sulfate electrode, equal to about −1.4 V versus Ag/AgCl (3.0 mol L−1 KCl)]. This electrode potential is actually not in the potential range of the hydrogen evolution reaction (see, for instance, Figure 5 in Mah´e, Devilliers, and Comninellis [19] or Figure 2 in Hupert et al. [54]); consequently, this mild pretreatment cannot be considered as a truly cathodic pretreatment in the sense used in previous studies (water electroreduction to produce hydrogen terminations) [18–20]. On the other hand, after the severe cathodic pretreatment {corresponding to a nearly stable electrode potential in the range of −3 V versus MSE [∼−3.4 V vs. Ag/AgCl (3.0 mol L−1 KCl)]}, the Ep value (700 mV with a small peak current increase) in the voltammograms registered with the ceric species (see Figure 8.8) is a bit smaller than that obtained for the as-deposited BDD electrode (750 mV), while Ep for Fe(CN)6 4−/3− redox system reaches the quasi-reversible value of 100 mV with a notable increase of the peak current [21]. In a subsequent study, Simon et al. [22] confirmed the enhancement in the conductivity of moderately doped (1019 cm−3 ) BDD surfaces after cathodic pretreatments of as-deposited samples. They also showed that previously annealed samples (1100◦ C in ultrahigh vacuum, for 12 h, in order to outgas the hydrogen introduced into the diamond layer during the deposition process) do not exhibit an enhancement of conductivity after the electrochemical pretreatments. This shows that the presence of hydrogen in the BDD electrodes seems to be crucial to increase the superficial conductivity after the electrochemical pretreatments.

190

CATHODIC TREATMENT OF BDD ELECTRODES AND THEIR USE IN ELECTROANALYSIS

400 (a) Ce3+/4+ in H2SO4 medium

j (µA cm–2)

200

0

NT –200

A1 C1

–400 –0.3

300

0.0

0.3

0.6 E (V/ESM)

0.9

1.2

1.5

(b) Fe(CN)64–/3– in KOH medium

j (µA cm–2)

200 100 0 –100

NT A1

–200

C1 –1.5

–1.0

–0.5 E (V/ESM)

0.0

0.5

Figure 8.7 Cyclic voltammograms (ν = 0.02 V s−1 ) for (a) 5.0 mmol L−1 Ce3+/4+ in 0.5 mol L−1 H2 SO4 and (b) 5.0 mmol L−1 Fe(CN)6 4−/3− in 0.5 mol L−1 KOH, at BDD electrodes after mild electrochemical pretreatments (±10−3 C cm−2 ): anodically (dashed lines) or cathodically (dotted lines) pretreated and as-deposited (full lines) electrodes. (Reprinted with permission from Ref. 21.)

A recent tentative to elucidate the electron transfer behavior on HT-BDD electrodes in an electrolytic solution was carried out by Wang et al. [55] using scanning probe microscopy and ab initio methods. The HT-BDD electronic structures were investigated in detail. Shallow acceptors were found in the HT-BDD bandgap and were attributed to the interaction between physisorbed active adsorbates and C–H bondings on the diamond surface. The authors also concluded that these shallow acceptors favor electron transfer. Studies of diamond electrodes in nonaqueous electrolytes are more limited. PastorMoreno and Riley [56] studied the reduction of 1,4-benzoquinone in acetonitrile at BDD electrodes with different surface pretreatments. They showed that the mechanism of reduction is dependent on the electrode pretreatment. Whereas the electrochemistry at an oxygenated BDD surface (prepared by immersion in a hot chromic acid solution) resembles that of a glassy-carbon electrode, the electrochemistry at an HT-BDD

8.2 CATHODIC PRETREATMENT OF CONDUCTIVE DIAMOND FILMS

400

191

(a) Ce3+/Ce4+ in H2SO4 medium

300

j (µA cm–2)

200 100 0 –100 NT –200

A2

–300

C2

–400 –0.5 600

0.0

0.5 E (V/ESM)

1.0

1.5

(b) Fe(CN)64–/3– in KOH medium

j (µA cm–2)

400 200 0 –200 NT

–400

A2 C2

–600 –1.0

–0.5 0.0 E (V/ESM)

0.5

Figure 8.8 Cyclic voltammograms (ν = 0.02 V s−1 ): (a) for 5.0 mmol L−1 Ce3+/4+ in 0.5 mol L−1 H2 SO4 and (b) for 5.0 mmol L−1 Fe(CN)6 4−/3− in 0.5 mol L−1 KOH, at BDD electrodes after severe electrochemical pretreatments (±1 C cm−2 ): anodically (dashed lines) or cathodically (dotted lines) pretreated and as-deposited (full lines) electrodes. (Reprinted with permission from Ref. 21.)

surface (as-prepared or obtained by applying a potential in the hydrogen evolution region, for 6 h, in 1 mol L−1 KCl) indicates the presence of a hydrogen source. In the latter case, the authors concluded that hydrogen in the diamond subsurface may participate in electrochemical processes. In view of the small number of reports available in the literature about the effect of cathodic treatments on the electrochemical and electroanalytical behavior and surface properties of BDD films, it is clear that this issue will demand further investigations in the future. For example, studies must be carried out to (1) elucidate the changes (in the bulk and in the surface) in the physical and/or physicochemical properties of diamond materials caused by cathodic pretreatments; (2) evaluate the physical and/or chemical stability of BDD films after repeated and severe cathodic pretreatments, focusing on the use of in situ

192

CATHODIC TREATMENT OF BDD ELECTRODES AND THEIR USE IN ELECTROANALYSIS

techniques; and (3) quantify and understand the relationship between different properties of diamond materials before and after cathodic pretreatments [properties such as sp2 type carbon content, grain boundary size (micro and nanodiamond), level of doping, conductivity, and surface termination, among others]. The results from these studies would certainly contribute to advance the fundamental knowledge and the technological applications of BDD in electrochemistry and electroanalysis. 8.3 8.3.1

ELECTROANALYTICAL APPLICATIONS General Aspects

One of the great challenges of analytical chemistry has been the fulfillment of the demand for analytical methods for the quantification of biological, environmental, and pharmaceutical analytes in several kinds of matrixes. The concentration of degradation products and/or impurities present in analytes varies widely and the analysis procedures usually require physical and/or chemical separation(s) prior to any meaningful determination. The assay of chemical species of interest may correspond to a very complex process, due to limiting factors that can be represented by the need to determine ever-smaller amounts (often below the LOD value offered by the techniques available), the interferences arising from the complex character of many matrixes, and the need to distinguish and quantify various chemical species associated with the same product (selectivity). As was pointed out previously in this chapter and by others (see, for instance, refs. 1–5), thin BDD films have an important number of electrochemical properties that distinguish them from other carbons bonded by sp2 , commonly used as electrodes in electroanalytical determinations. Because of these properties, BDD electrodes have been widely studied in recent years, both in terms of fundamental electrochemical properties [1–3,18,20,35,57,58] and electroanalytical [1,4–6,59,60] and environmental (wastewater treatment) applications [61–67]. The wide potential window (up to 3.5 V) presented by BDD electrodes in aqueous solutions—as well as their very low and stable voltammetric background current, long-term response stability, and low sensitivity to dissolved oxygen—are important characteristics that allow the detection of many electroactive species, which otherwise would be masked by the water decomposition reactions. However, it is clear that the analytical performance of BDD electrodes greatly depends on their surface termination (i.e., whether they are hydrogen or oxygen terminated [68–74]). Accordingly, surface modifications of BDD have a strong effect on its electroanalytical behavior when detecting several kinds of analytes. The presence of different functional groups on the BDD surface plays an important role when using diamond electrodes in electroanalysis. Initially, in most studies found in the literature, diamond electrodes were subjected mainly to an anodic treatment of the surface in order to make it hydrophilic (standard procedure) [74–76]. However, as pointed out in Section 8.2, more recently Suffredini et al. [18] and Salazar-Banda et al. [20] highlighted that cathodic pretreatments of BDD electrodes lead to a nearly-ideal reversible behavior for the Fe(CN)6 4−/3− redox system, along with improving the electrochemical responses for other analytes in aqueous media. So, this pretreatment has been used in many studies, especially those involving measurements for the detection of pesticides, drugs, and food additives using classical voltammetric techniques [e.g., square-wave voltammetry (SWV) and differential pulse voltammetry (DPV)]. It has been shown that for many analytes the combination

8.3 ELECTROANALYTICAL APPLICATIONS

193

of a cathodically pretreated (hydrogen-terminated) BDD electrode with these techniques becomes a very powerful analytical tool. Hence, a summary of electroanalytical applications of cathodically pretreated BDD electrodes in the determination of several kinds of analytes in different matrixes is presented hereinafter. 8.3.2

Determination of Pesticides in Environmental Samples

Although chemical substances such as herbicides, germicides, and insecticides (pesticides) are necessary to improve agricultural productivity, they can cause prejudicial effects on the environment since they flow into natural or public water systems. Commonly, most of them are carcinogenic; consequently, their presence in the environment can also be harmful to human health. Since most electroanalytical studies to determine pesticides are done using a mercury electrode, the search for new materials becomes increasingly important. The electroanalytical determination of some pesticides using cathodically pretreated BDD electrodes is summarized here. 8.3.2.1 Carbaryl Carbaryl, like any other insecticide, must be toxic to insects to be effective; like all pesticides, it is also toxic to certain nontarget organisms, including humans. Carbaryl acts both by entering the stomach of the pest with food and by being absorbed through body contact. The occupational exposure of humans to this insecticide has been observed to cause cholinesterase inhibition and reduction of the activity of this enzyme in the blood, which may cause neurological effects [77]. Therefore, simple, sensitive, and reliable methods for the analysis of carbaryl in air, food, grains, and natural water are quite welcome. Codognoto et al. [78] reported the development of an analytical method to determine carbaryl, in aqueous solutions as well as in natural waters, using a BDD electrode and SWV. Prior to the determinations, the electrode was first anodically pretreated [+ 3.0 V versus Ag/AgCl (3.0 mol L−1 KCl) for 30 min] to clean its surface, followed by a cathodic pretreatment [−3.0 V versus Ag/AgCl (3.0 mol L−1 KCl) for 30 min]. The developed method of analysis does not involve any previous extraction, cleanup, preconcentration, or derivatization steps. The obtained analytical curve range was 2.5–30.0 µmol L−1 (see Figure 8.9), with a limit of detection (LOD) of 8.2 ± 0.2 µg L−1 in pure water (analytical sensitivity of 3.07 mA mmol−1 L) and a limit of quantification (LOQ) of 27.5 µg L−1 ; the attained reproducibility and repeatability (n = 10) were 3.9% and 3.2%, respectively. This analytical sensitivity was slightly decreased (to 2.80–2.90 mA mmol−1 L) when the experiments were carried out using water samples collected from two different points in a polluted urban creek. The obtained LOD value is quite adequate, considering that according to the local legislation the maximum amount of carbaryl allowed in natural water is 10 µg L−1 . The authors also evaluated the effect of other pesticides (fenthion and 4-nitrophenol) on the carbaryl determination and found an insignificant influence, as shown in Figure 8.10. In other words, the voltammetric procedure reported by the authors leads to clearly separated electrochemical responses for the three analytes, which can thus be determined simultaneously. 8.3.2.2 4-Nitrophenol 4-Nitrophenol (4-NP), a hazardous substance that can have a high environmental impact due to its toxicity and persistence, is a metabolite of several

194

CATHODIC TREATMENT OF BDD ELECTRODES AND THEIR USE IN ELECTROANALYSIS

Figure 8.9 Calibration curve from SWV responses of a cathodically pretreated BDD electrode (A = 0.62 cm2 ) for different carbaryl concentrations in the range 2.5–30.0 μmol L−1 in 0.1 mol L−1 Na2 SO4 (pH 6.0). SWV conditions: Es = 2 mV, a = 50 mV, f = 300 Hz. Inset: linear dependence of the peak current on the carbaryl concentration. (Reprinted from Ref. 78.)

Figure 8.10 SWV profile on the BDD electrode (A = 0.62 cm2 ) for 10 μmol L−1 4-nitrophenol (1), 10 μmol L−1 fenthion (2), and 25 μmol L−1 carbaryl (3) in a 0.1 mol L−1 Na2 SO4 (pH 6.0) solution. SWV conditions: Es = 2 mV, a = 50 mV, f = 300 Hz. (Reprinted from Ref. 78.)

organophosphorus pesticides. Thus, some electroanalytical methods have been developed to determine its residues in different types of samples. Pedrosa, Codognoto, and Avaca [79] reported the electroanalytical determination of 4-NP in aqueous solutions on a BDD electrode using SWV. Prior to the experiments the electrode was cathodically pretreated at − 3.0 V versus Ag/AgCl (3.0 mol L−1 KCl) for 30 s. The compound presented only one irreversible peak for both its oxidation at 1.0 V versus Ag/AgCl (3.0 mol L−1 KCl) and its reduction at −0.8 V versus

8.3 ELECTROANALYTICAL APPLICATIONS

195

Ag/AgCl (3.0 mol L−1 KCl), as seen in Figure 8.11. The corresponding LOD values are 8.4 µg L−1 and 12.1 µg L−1 , respectively. Garbellini, Salazar-Banda, and Avaca [80] also investigated the determination of 4-NP on a cathodically pretreated BDD electrode using SWV, but associated with ultrasound radiation. Since this radiation can clean the electrode surface, its fouling commonly caused by strong adsorption of the electroactive species could be minimized and, consequently, the LOD values were improved. Prior to the experiments, the BDD electrode was pretreated in a 0.5 mol L−1 H2 SO4 solution by applying + 3.0 V versus Ag/AgCl (3.0 mol L−1 KCl) for 5 s, followed by −3.0 V Ag/AgCl (3.0 mol L−1 KCl) for 30 s. As can be apprehended when Figure 8.12 is compared with Figure 8.11, clearly the application of ultrasound radiation increased the sensibility of the detection method compared to the silent conditions (about 3 times higher). This was explained by the authors as due 20

20

Ip/µA

8

10 5

15 I (µA)

0

(a)

9

15

7 0

5 1 2 3 4 –5 –1 [4–Nitrofenol]/10 mol L

6

10 5 4 3 2 1

5

0 0.8

0.9

1.0

1.2

1.1

E (V) vs Ag/AgCI 15

–15

9

(b)

–12

–Ip/µA

12 9

8

6

I (µA)

3

–9

0

7 0

1 2 3 [4–Nitrofenol] / 10–5mol L–1

6 5

–6 4 3

–3

2 1

0 –0.4

–0.5

–0.6

–0.7

–0.8

–0.9

–1.0

–1.1

E (V) vs Ag/AgCI

Figure 8.11 SWV response of the BDD electrode (A = 0.62 cm2 ) for different 4-NP concentrations in a 0.1 mol L−1 Britton-Robson buffer (pH 6.0). (a) oxidation and (b) reduction-based determinations. SWV conditions: a = 60 mV, Es = 2 mV, f = 100 Hz. Insets: linear dependence of the peak current with 4-NP concentration. (Reprinted with permission from Ref. 79.)

196

CATHODIC TREATMENT OF BDD ELECTRODES AND THEIR USE IN ELECTROANALYSIS

50 40 Ip/µA

30

12

30

11

20

10

A

I (µA)

10

9

0

20

(a)

13

B

40

8 0

10

20 30 40 50 –6 –1 [4–NP]/10 mol L

60 7 6

10

5 4 1

0 0.6

0.8

3 2

1.0

1.2

E (V) vs Ag/AgCI –120

13

−140

12

−100

11

Ip/µA

–100 –80

B

−120 −80 −60

10

A

−40

9

I (µA)

−20 0

–60

0

10

20

30

[4–NP]/10

–6

40

50

60

8

–1

mol L

–40

7 6 5 4

–20

3 2 1

0 –0.2

(b)

–0.4

–0.6

–0.8

–1.0

E (V) vs Ag/AgCI

Figure 8.12 SWV responses obtained on a BDD electrode (A = 0.25 cm2 ) using different 4-NP concentrations in a 0.1 mol L−1 Britton-Robson buffer (pH = 6.0) in the presence of ultrasound radiation. (a) oxidation and (b) reduction-based determinations. SWV conditions: f = 100 Hz, a = 50 mV, Es = 2 mV. Ultrasound conditions: d: 5 mm and A: 20%. Insets: The respective analytical curves for the oxidation and reduction processes under silent (A) and ultrasound conditions (B). (Reprinted with permission from Ref. 80.)

to the increase in mass transport and the cleaning of the electrode surface brought on by the ultrasound radiation. Furthermore, the sensitivity of the method based on the 4-NP reduction process was increased much more significantly than that based on the oxidation process. The obtained LOD values were 3.87 µg L−1 for the oxidation process and 2.57 µg L−1 for the reduction process; the corresponding LOQ values were 12.9 µg L−1 and 8.58 µg L−1 , respectively. 8.3.2.3 Chlorophenols Pentachlorophenol (PCP) is a very toxic compound and among the chlorophenols it was extensively used for decades as a fungicide in wood preservation. Its analysis has received special attention in order to measure the amounts of this substance that can be transferred to food by direct contact with wood treated with

8.3 ELECTROANALYTICAL APPLICATIONS

197

PCP. Besides, it can be used as model compound for the development of new analytical techniques [43]. The toxicity and persistence of PCP in contaminated places is well known, so that it is considered as a toxicologically relevant pollutant [82]. Codognoto, Machado, and Avaca [43] reported on the optimization of the experimental parameters for the determination of PCP in pure and contaminated water on a BDD electrode also using SWV. Prior to the experiments, the electrode was cathodically pretreated by applying −3.0 V versus Ag/AgCl (3.0 mol L−1 KCl) for 30 s. The obtained calibration curves for the determination of PCP showed an excellent linear response, for solutions both in pure and contaminated water, but with different slopes (see Figure 8.13). Possible explanations for this behavior, which leads to different LOD values, were put forth by the authors: hindrance of a 4-CP adsorption step by contaminant organic molecules, decrease of the concentration of 4-CP or other chlorophenols by interaction with humic and fulvic substances, or presence of other unknown species that undergo oxidation in the same potential region. The optimized method yielded LOD values for PCP of 5.5 mg L−1 and 15.5 mg L−1 in pure and contaminated water, respectively. According to the authors, the LOD value obtained in pure water was low enough to comply with the limits imposed by environmental and/or health authorities for human consumption, whereas that obtained in polluted matrixes was just above that limit. Two years later, Suffredini et al. [18] clearly demonstrated that the electrochemical response of a BDD electrode in the determination of 4-chlorophenol or PCP was strongly affected by the type of pretreatment applied to its surface before measurements (see Section 8.2). Pedrosa, Machado, and Avaca [45] reported on the simultaneous determination of 4chlorophenol (4-CP), 2,4-dichlorophenol (2,4-CP), and 2,4,6-trichlorophenol (2,4,6-CP) in different water samples on a cathodically pretreated BDD electrode through a deconvolution procedure. This procedure was used in order to obtain a separation of the peaks

Figure 8.13 Analytical curves for PCP in solutions prepared with pure water () and with samples collected at point 1 (•), point 2 (⋄), and point 3 () of a contaminated urban creek. SWV conditions: a = 60 mV, Es = 2 mV, f = 100 Hz. BDD electrode area = 0.62 cm2 . (Reprinted with permission from Ref. 43.)

198

CATHODIC TREATMENT OF BDD ELECTRODES AND THEIR USE IN ELECTROANALYSIS

TABLE 8.1 Analytical parameters for 4-CP determination in pure and contaminated water (from two different points of a polluted creek) samples, using SWV with a cathodically pretreated BDD electrode or an HPLC method. (Reprinted with permission from Ref. 45.) Method BDD HPLC

Sample

R

LOQ (µg L−1 )

Recovery (%)

Milli Q Point 1 Point 2 Milli Q Point 1 Point 2

0.998 0.995 0.996 0.999 0.998 0.998

9.2 37.2 42.2 1.2 36.1 38.4

94.0 ± 3.1 92.2 ± 4.2 91.1 ± 3.5 96.0 ± 1.1 94.2 ± 5.2 95.1 ± 4.8

for each compound and to construct standard curves for each chlorophenol in the mixture. Analytical curves were obtained with the cathodically pretreated BDD electrode by the standard addition method for the mixture of chlorophenols in the different water samples (pure or collected in two different points of a polluted urban creek) and compared with those obtained using an HPLC method. Recovery experiments were also carried out. From the obtained analytical results, summarized in Table 8.1, the authors concluded that excellent recoveries were attainable even in highly polluted water samples, thus indicating that the reported procedure could be used for determinations in environmental matrices. A slight dependence of the LOQ values on the amount of pollution in the water samples was attributed to the same factors put forth by Codognoto, Machado, and Avaca [43] to explain the variation of the slope of the analytical curves for PCP in pure and contaminated water.

8.3.3

Determination of Substances in Food Samples

8.3.3.1 Aspartame Aspartame (N-L-α-aspartyl-L-phenylalanine methyl ester) has the potential to reduce the amount of sugars and calories in food products, and thus may be combined with sugars (such as sucrose, dextrose, and fructose) and/or sweeteners (such as acesulfame-K, sodium cyclamate, and saccharin). However, aspartame is not always suitable for baking because it often breaks down when heated and it loses much of its sweetness. Considering that most methods developed for aspartame determination require a lengthy pretreatment of the sample prior to analyses, Medeiros et al. [68] proposed an electroanalytical method using a BDD electrode and SWV to determine the aspartame concentration directly in the sample without pretreatment or chemical separation. Prior to the experiments, the BDD electrode was cathodically pretreated by applying −1.0 A cm−2 for 60 s in a 0.5 mol L−1 H2 SO4 solution. As shown in Figure 8.14, the oxidation of aspartame on BDD presents a single peak at 1.6 V versus Ag/AgCl (3.0 mol L−1 KCl), with the characteristics of an irreversible reaction. The obtained analytical curve is linear in the aspartame concentration range 9.9–52 µmol L−1 , with an LOD of 0.23 µmol L−1 ; in repeatability studies (n = 5) of a 0.10 mmol L−1 aspartame solution, the obtained relative standard deviation (RSD) was 0.2%. Several dietary products containing aspartame were analyzed using the proposed SWV method and employing the standard addition method. The inset in Figure 8.14b is the obtained analytical curve

8.3 ELECTROANALYTICAL APPLICATIONS

12 R = 0.9998 10 8 6 4 2 0 0 1 2 3 4 5 6 [Aspartame]/10–5 mol L–1

11 10 9 8 7 6 5 4 3 2

I (µA)

25

I (µA)

20

199

15

10

1

5

(a) 0 1.3

1.4

1.5

1.6

1.7

1.8

1.9

E (V) vs. Ag/AgCl 14 12 10 8 6 4 2 0 –3 –2 –1 0 1 2 3 4 [Aspartame]/10–5 mol L–1

5

I (µA)

25

20

4 3 2

15 I (µA)

1 10

5 (b)

0 1.3

1.4

1.5

1.6

1.7

1.8

E (V) vs. Ag/AgCl

Figure 8.14 (a) SWV response of a cathodically pretreated BDD electrode (A = 0.72 cm2 ) for the oxidation of aspartame in increasing concentrations [9.9–52 μmol L−1 , in 0.5 mol L−1 H2 SO4 (curve 1)]; Inset: corresponding analytical curve. (b) Square-wave voltammograms obtained for the determination of aspartame in a sweetener (Zero cal®). SWV conditions: a = 40 mV, Es = 2 mV, f = 10 Hz. (Reprinted with permission from Ref. 68.)

for the determination of aspartame in a sample of a dietary product, the sweetener Zero cal®. 8.3.3.2 Sodium Cyclamate Sodium cyclamate is a white, crystalline, and odorless powder. Although 30 times sweeter than sucrose, it is noncaloric (zero calories) and more stable than other artificial sweeteners, such as aspartame and saccharin, which allows its use at high and low temperatures [82]. Sodium cyclamate is commonly used together with saccharin because cyclamate can mask the bitter flavor left by saccharin, but its

200

CATHODIC TREATMENT OF BDD ELECTRODES AND THEIR USE IN ELECTROANALYSIS

use has been banned in the US because it was found that it could be metabolized as cyclohexylamine, which can cause several health problems. So, taking into account the interest in the development of analytical methods for the determination of these sweeteners in dietary products, Medeiros et al. [69] proposed the use of SWV with a BDD electrode for the analytical determination of sodium cyclamate. The BDD electrode was cathodically pretreated the same way as it was done for the determination of aspartame [68] (discussed earlier). Figure 8.15 clearly shows the enhancement of the BDD electrochemical activity brought on by this cathodic pretreatment, compared to its untreated (as-received) response. The samples were analyzed as received in a 0.5 mol L−1 H2 SO4 solution in the concentration range 5.0 × 10−5 −4.1 × 10−4 mol L−1 , with an LOD of 4.8 × 10−6 mol L−1 . In the repeatability and reproducibility studies (n = 5) of a 3.0 mmol L−1 cyclamate solution, RSD values of 1.2% and 2.4% were obtained, respectively; furthermore, the proposed method was applied with success in the determination of sodium cyclamate in several dietary products. 8.3.3.3 Aspartame and Sodium Cyclamate Medeiros et al. [70] also reported on the simultaneous determination of aspartame and sodium cyclamate in dietary products using SWV and a cathodically pretreated BDD electrode. The SWV oxidation peak potentials of aspartame and cyclamate present in binary mixtures are about 400 mV apart, as shown in Figure 8.16. For aspartame, LOD was 0.47 µmol L−1 in the presence of 0.30 mmol L−1 cyclamate; for sodium cyclamate, LOD was 4.2 µmol L−1 in the presence of 0.10 mmol L−1 aspartame. When simultaneously changing the concentration of both sweeteners in a 0.5 mol L−1 H2 SO4 solution, the corresponding LOD values were 0.35 and 4.5 µmol L−1 , respectively. In repeatability tests (n = 5), the obtained RSD values were 1.3%, for a 0.10 mmol L−1 aspartame solution, and 1.1% for a 3.0 mmol L−1 cyclamate solution. The proposed voltammetric method

200 a b

I (µA)

150

100

50

0 1.6

1.7

1.8

1.9

2.0

2.1

E (V) vs. Ag/AgCl (KCl 3.0 mol L–1)

Figure 8.15 SWV responses for 3.0 mmol L−1 sodium cyclamate in 0.5 mol L−1 H2 SO4 using an as-received (a) or a cathodically pretreated (b) BDD electrode. SWV conditions: a = 20 mV, Es = 2 mV, f = 10 Hz. BDD electrode area = 0.72 cm2 . (Reprinted with permission from Ref. 69.)

8.3 ELECTROANALYTICAL APPLICATIONS

201

(b) (a)

(c)

Figure 8.16 (a) SW voltammograms obtained for the oxidation of aspartame and cyclamate in 0.5 mol L−1 H2 SO4 . The concentrations of both aspartame (5.0–50 μmol L−1 ) and cyclamate (0.050–0.50 mmol L−1 ) were changed simultaneously. (b) Analytical curves for aspartame. (c) Analytical curves for cyclamate. SWV conditions: a = 40 mV, Es = 2 mV, f = 10 Hz. BDD electrode area = 0.72 cm2 . (Reprinted with permission from Ref. 70.)

was successfully applied in the simultaneous determination of aspartame and sodium cyclamate in several dietary products, with results in very good agreement with those results obtained using an HPLC method. 8.3.3.4 Total Phenols Total phenol concentration is an essential indicator of the state of quality of pharmaceutical products and food, since its concentration could be related to nutritional, processing, and health aspects. Thus, Dejmkova et al. [83] report on a method to determine total phenols in foods such as teas, juices, and wines. The electrode was electrochemically pretreated in 1 mol L−1 HNO3 by a procedure similar to the one proposed by Suffredini et al. [18]: +3.0 V versus Ag/AgCl (3 mol L−1 KCl) for 20 min, followed by −3.0 V versus Ag/AgCl (3 mol L−1 KCl) also for 20 min. When the applied potential on the cathodically pretreated BDD is positive enough to oxidize the phenols present in the sample, the charge consumed during the oxidation could be related to the total phenol content in the sample. The proposed method is robust, allowing a rapid evaluation of total phenols directly in real samples after a simple dilution with the supporting electrolyte. The electrode does not present the drawback of fouling (an electrochemical cleaning procedure was successfully optimized) and the results were validated using the standard Folin-Ciocalteau assay. 8.3.4

Determination of Substances in Pharmaceutical Samples

8.3.4.1 Sulfamethoxazole and Trimethoprim Sulfonamides, indicated primarily to treat urinary infections, are used in combination with trimethoprim (TMP) for the treatment of ear infections, bronchitis, sinusitis, and pneumocystis pneumonia. The corresponding pharmaceutical products usually consist of a sulfonamide mixed with another

202

CATHODIC TREATMENT OF BDD ELECTRODES AND THEIR USE IN ELECTROANALYSIS

µ

drug that increases its power, such as the sulfamethoxazole (SMX) and TMP mixture, an often used pharmaceutical product. In the specific case of analysis of drugs, quality control can be considered as one of the most important aspects because it contributes to ensure the efficacy, safety, and fundamentally the quality standards of medicines. Thus, the search for techniques that are sensitive, accurate, and easily accessible for the determination of compounds present in several commercial drugs has received much attention. In this context, Andrade et al. [73] reported on the simultaneous electrochemical detection by differential pulse voltammetry (DPV) of SMX and TMP using cathodically (HT-) or anodically (OT-) pretreated BDD electrodes. The cathodic or anodic pretreatment was carried out by applying −0.5 A cm−2 or 0.5 A cm−2 , respectively, during 60 s, in a 0.5 mol L−1 H2 SO4 solution. Cyclic voltammetric studies (see Figure 8.17) show that on an HT-BDD electrode both the SMX and TMP voltammograms exhibit well-defined irreversible oxidation peaks at 1080 mV and 1100 mV versus Ag/AgCl (3.0 mol L−1 KCl), respectively. Conversely, within the investigated potential range (0.5 V to 1.3 V versus Ag/AgCl), no reduction peaks are observed on the reverse scan. However, since the anodic peak potentials for SMX and TMP are separated by only 20 mV, their simultaneous analyses would not be possible under these conditions. Thus, by changing the solution pH, a greatly increased difference in the oxidation peak potential values (180 mV) was found at pH 7 (see Figure 8.18a), associated to two well-defined oxidation waves [Ep = 920 and 1100 mV versus Ag/AgCl (3.0 mol L−1 KCl)] that correspond to the oxidation of SMX and TMP, respectively. When an OT-BDD electrode was used, the magnitude of these oxidation waves decreased, but this decrease was more significant for the SMX oxidation. As can be seen in Figure 8.18b, independently of whether an HT-BDD or an OT-BDD electrode is used, the SMX oxidation current peak decreases as the pH increases. Besides, the magnitude of the current peak obtained for the SMX and TMP oxidation was always higher when the HT-BDD electrode was used.

Figure 8.17 Cyclic voltammograms (ν = 50 mV s−1 ) for (a) blank solution (0.2 mol L−1 Britton-Robison buffer, pH 2.0) and solutions containing (b) 1.0 mg L−1 SMX or (c) 1.0 mg L−1 TMP in this buffer, on a cathodically pretreated BDD electrode. BDD electrode area = 0.63 cm2 . (Reprinted from Ref. 73.)

8.3 ELECTROANALYTICAL APPLICATIONS

203

(a)

(b)

Figure 8.18 (a) Differential pulse voltammetric responses (oxidation) obtained at an anodically (solid line) or cathodically (dashed line) pretreated BDD electrode (A = 0.63 cm2 ) using a mixture of 1.0 mg L−1 SMX and 1.0 mg L−1 TMP in a 0.2 mol L−1 Britton-Robson buffer (pH 7.0). (b) Effect of pH on the SMX oxidation peak current at an anodically (solid line) or cathodically (dashed line) pretreated BDD electrode. DPV conditions: scan rate, 50 mV s−1 ; pulse amplitude, 60 mV; pulse width, 10 ms. (Reprinted from Ref. 73.)

In an effort to explain the different electrochemical behaviors obtained when using an HT- or an OT-BDD electrode, Andrade et al. [73] analyzed the acid-base chemistry of SMX and TMP. The former is a weak acid (pKa = 5.6), whereas the latter is a weak base (pKa = 7.3). Thus, if pH > pKa , the conjugate base predominates in the solution bulk, the analyte then being negatively charged. Consequently, an electrostatic repulsion between sulfa anions and the OT-BDD electrode could be expected for the measurements carried out at pH > 5.6. In fact, for pH > 5.6, the current peaks obtained for the SMX oxidation on the OT-BDD electrode (see Figure 8.18b) were always lower than those obtained on the HT-BDD electrode. However, when the measurements were

204

CATHODIC TREATMENT OF BDD ELECTRODES AND THEIR USE IN ELECTROANALYSIS

done at pH < 5.6, the SMX oxidation current peaks were not higher on the OT-BDD electrode when compared with those obtained on the HT-BDD electrode, as it could be expected. Furthermore, in the pH range investigated (2 to 7) the oxidation current peaks always decreased linearly with pH, independently of whether an HT-BDD or an OT-BDD electrode was used. Conversely, for pH < 7.3, the magnitude of the TMP oxidation current peak could be expected to be higher on the OT-BDD electrode than on the HT-BDD electrode, but actually it was about 40% lower. Taking this into account, the authors concluded that the electroanalytical performance of the electrodes is not significantly influenced by electrostatic interaction between the analytes and the BDD surface; hence, the obtained results should be explained mainly taking into account the different surface conductivity presented by the HT-BDD or OT-BDD electrodes. However, Ivandini et al., [84] have pointed out that while HT-BDD acts as an excellent electrode for the detection of negatively charged DNA, OT-BDD works exceptionally well for the detection of the positively charged oxidized form of glutathione [85], due to the operation of strong electrostatic interactions. Clearly, questions involving the influence of surface chemistry of the BDD films on the electrochemical processes occurring on these electrodes are still controversial, needing to be further investigated. A series of DPV voltammograms obtained for the simultaneous determination of SMX and TMP at different concentrations (1.0−10 mg L−1 and 0.2−2.0 mg L−1 for SMX and TMP, respectively) in a 0.2 mol L−1 Britton-Robinson buffer (pH 7) by DPV at an HT-BDD electrode are shown in Figure 8.19. Notice in the figure insets that the respective analytical curves presented a good linearity in the investigated concentration range (r 2 = 0.9993 for both SMX and TMP). The calculated values for LOD and LOQ were 3.65 µg L−1 (14.4 nmol L−1 ) and 12.2 µg L−1 (48.2 nmol L−1 ) for SMX; these values were 3.92 µg L−1 (13.5 nmol L−1 ) and 13.1 µg L−1 (45.1 nmol L−1 ) for TMP.

Figure 8.19 Differential pulse voltammetric responses (simultaneous oxidation) obtained on an HTBDD electrode (A = 0.63 cm2 ) at different concentrations of SMX/TMP in a 0.2 mol L−1 Britton-Robson buffer solution (pH 7.0). SMX/TMP concentrations: (1) 1.0/0.2; (2) 2.0/0.4; (3) 3.0/0.6; (4) 4.0/0.8; (5) 5.0/1.0; (6) 6.0/1.2; (7) 7.0/1.4; (8) 8.0/1.6; (9) 9.0/1.8, and (10) 10/2.0 mg L−1 . Inset: respective analytical curves for SMX and TMP. DPV conditions: scan rate, 50 mV s−1 ; pulse amplitude, 60 mV; pulse width, 10 ms. (Reprinted from Ref. 73.)

8.3 ELECTROANALYTICAL APPLICATIONS

205

Besides, repeatability tests carried out by successive measurements (n = 10) in the same solution (10 mg L−1 SMX and 2.0 mg L−1 TMP) showed RSD values of 0.3% and 0.1%, respectively. Despite the fact that LOD and LOQ are not considered so relevant in drugs determination, because of their high concentration in commercial formulations, the obtained values clearly indicate that quite low concentrations of SMX and TMP can be detected using the HT-BDD electrode. So, the proposed method was applied successfully to determine SMX and TMP by the standard addition method in three different commercial formulations. 8.3.4.2 Sulfamethoxazole and Sulfadiazine Souza et al. [86] also reported on the electrochemical determination of sulfonamides (sulfadiazine and sulfamethoxazole, independently) in pharmaceutical formulations employing a cathodically pretreated BDD electrode and SWV. The BDD electrode was pretreated in 0.5 mol L−1 H2 SO4 similarly to what was done by Salazar-Banda et al. [20]. The as-received electrode was first anodically pretreated (+3.0 V versus SCE for 30 min) to clean its surface, followed by a cathodic pretreatment [−3.0 V versus SCE for 30 min). Then, before each measurement, this cathodic pretreatment was carried out for 30 s (only for the very first pretreatment each day, the applied potential was −2.0 V versus SCE). Good linear analytical curves were achieved for both sulfonamides and the obtained LOD values were 2.19 and 1.15 µmol L−1 for sulfadiazine and SMX, respectively. In the repeatability studies (n = 7) of 90 µmol L−1 sulfadiazine and 0.25 mmol L−1 SMX solutions, RSD values of 0.56% and 0.58% were obtained, respectively. In the corresponding reproducibility study (n = 5), RSD values of 0.78% and 0.71% were obtained for sulfadiazine and SMX, respectively. The recovery values for both sulfonamides were in the range 95–104% and the methodology was successfully compared with the standard HPLC method, with relative errors of −4.31% and −0.79% for sulfadiazine and sulfamethoxazole, respectively. 8.3.4.3 Acetylsalicylic Acid Acetylsalicylic acid (ASA), also known by the trade name aspirin, is one of the oldest medicines that still play an important role in modern therapeutics, being widely employed in pharmaceutical formulations for the relief of headaches, fever, muscular pain, and inflammations caused by arthritis or injury. Commonly, ASA is determined by titration after its conversion to salicylic acid and acetic acid by alkaline hydrolysis. Aiming at obtaining a simpler procedure, Sartori et al. [87] investigated the determination of ASA in pharmaceutical formulations using SWV and a cathodically pretreated BDD electrode. In this proposed electroanalytical method, ASA can be directly determined in a 0.01 mol L−1 H2 SO4 solution (see Figure 8.20). The obtained analytical curve was linear in the ASA concentration range 2.50 × 10−6 −1.05 × 10− mol L−1 , with a LOD of 2.0 µmol L−1 ; in the repeatability study (n = 10) of 45 µmol L−1 ASA solutions, a RSD value of 1.4% was obtained. The proposed method was applied with success in the determination of ASA in several pharmaceutical formulations (commercial adult and children tablets) and the obtained results were in close agreement, at a 95% confidence level, with those obtained using an official method of the British Pharmacopoeia. The reported results demonstrate that the combination of SWV with an HT-BDD electrode is a feasible alternative for the analytical determination of ASA in commercial adult and children tablets without previous hydrolysis of the analyte.

206

CATHODIC TREATMENT OF BDD ELECTRODES AND THEIR USE IN ELECTROANALYSIS

50

50 I /µA

40

40

30 20 10

I /µA

0 0 20 40 60 80 100 120 –6 –1 [AAS]/10 mol L

30 20

17 10 1 0 1.7

1.8

1.9 2.0 E (V) vs Ag/AgCl

2.1

2.2

Figure 8.20 SWV response (direct current) of an HT-BDD electrode (A = 0.33 cm2 ) for different ASA concentrations in 0.01 mol L−1 H2 SO4 : (1) 0; (2) 2.50 × 10−6 ; (3) 5.00 × 10−6 ; (4) 7.50 × 10−6 ; (5) 1.50 × 10−5 ; (6) 2.25 × 10−5 ; (7) 3.00 × 10−5 ; (8) 3.75 × 10−5 ; (9) 4.50 × 10−5 ; (10) 5.25 × 10−5 ; (11) 6.00 × 10−5 ; (12) 6.75 × 10−5 ; (13) 7.50 × 10−5 ; (14) 8.25 × 10−5 ; (15) 9.00 × 10−5 ; (16) 9.75 × 10−5 ; (17) 1.05 × 10−4 mol L−1 . Inset: analytical curve for the ASA oxidation process. SWV conditions: a = 40 mV, Es = 3 mV, f = 50 Hz. (Reprinted with permission from Ref. 87.)

8.3.4.4 Paracetamol and Caffeine Paracetamol (N-acetyl-p-aminophenol, acetaminophen), a substance derived from p-aminophenol, is considered an excellent analgesic in cases of mild to moderate pain, besides not causing relevant gastrointestinal effects. It is noncarcinogenic and an effective substitute for aspirin for patients with sensitivity to it. Caffeine is used therapeutically in combination with ergotamine in the treatment of migraines or with nonsteroidal anti-inflamatories in analgesic formulations. Lourenc¸a˜ o et al. [88] investigated the use of a cathodically pretreated BDD electrode to develop simple, selective, and sensitive methods for the determination of paracetamol and caffeine, simultaneously and individually. This was achieved using (1) SWV, for paracetamol; and (2) DPV, for caffeine individually and for both drugs simultaneously. The HT-BDD electrode was obtained by applying −1.0 A cm−2 for 180 s in a 0.5 mol L−1 H2 SO4 solution. In the binary mixtures, a separation of about 550 mV between the peak oxidation potentials of paracetamol and caffeine was obtained. Figure 8.21 shows the DPV voltammograms obtained for the simultaneous determination of the drugs. The corresponding calibration curves showed an excellent linear response, in the range 0.50–83 µmol L−1 for both compounds. The LOD values for the simultaneous determination of paracetamol and caffeine were 0.49 and 0.035 µmol L−1 , respectively. In the repeatability study (n = 5) of a 50 µmol L−1 paracetamol and caffeine solution, RSD values of 0.3% and 1.8% were obtained, respectively. The proposed method was successfully applied in the simultaneous determination of paracetamol and caffeine in several pharmaceutical formulations (tablets), with results similar to those obtained using a reference HPLC method. 8.3.4.5 Sildenafil Citrate (Viagra) Sildenafil citrate (1-[[3-(6,7-dihydro-1methyl-7-oxo-3-propyl-1-H-pyrazolo[4,3-d]pirydin-5-yl)-4-ethoxyphenyl]sulfonyl]-4methylpiperazine citrate), commonly known as Viagra®, is a drug widely used as oral

8.3 ELECTROANALYTICAL APPLICATIONS

207

200

I (µA)

150

100 15 50 1 0 0.4

0.6

0.8 1.0 1.2 E (V) vs. Ag/AgCl

1.4

1.6

Figure 8.21 Differential pulse voltammetric curves obtained for the oxidation of paracetamol and caffeine at equal concentrations in a 0.2 mol L−1 acetate buffer solution (pH 4.5): (1) 0.50, (2) 2.0, (3) 4.0, (4) 5.9, (5) 7.9, (6) 9.8, (7) 19, (8) 28, (9) 37, (10) 45, (11) 54, (12) 61, (13) 69, (14) 76, and (15) 83 μmol L−1 . DPV conditions: scan rate, 70 mV s−1 ; modulation amplitude, 100 mV; modulation time, 7 ms. BDD electrode area = 0.72 cm2 . (Reprinted with permission from Ref. 88.)

therapy for erectile dysfunction. Due to the high consumption of sildenafil, selective, sensitive, and easily used methods of analysis of its pharmaceutical formulations are welcome, to assess not only their quality but also their authenticity. Accordingly, Batista et al. [89] proposed the use of DPV in conjunction with a cathodically (−1.0 A cm−2 for 240 s in 0.5 mol L−1 H2 SO4 ) pretreated BDD electrode for the analytical determination of commercial pharmaceutical samples of Viagra®. According to the authors, the HT-BDD electrode presented a better peak definition and a higher current magnitude, indicating that the cathodic pretreatment of the electrode led to a larger electrochemical activity for sildenafil oxidation. The cyclic voltammetric response of sidenafil presents two electrochemically irreversible anodic peaks, at ∼1.5 and ∼2.0 V versus Ag/AgCl (3.0 mol L−1 KCl). In order to avoid interference from the oxygen evolution reaction, only the first peak was considered for the development of the electroanalytical method. After optimization of the DPV parameters, which included the maximum peak current and the minimum half-peak width values, the authors validated the method taking into account selectivity (possible interferents), linearity, and recovery studies. All these validation criteria were satisfied and the obtained LOD value was 0.64 µmol L−1 . In the repeatability (n = 10) and reproducibility (n = 5) studies of 4.0 µmol L−1 sildenafil solutions, RSD values of 1.1% and 1.9% were obtained, respectively. Table 8.2 summarizes the results obtained for determinations of sildenafil citrate in Viagra® commercial tablets of different dosages employing the proposed DPV method compared with those obtained using a reference HPLC method. 8.3.4.6 Lidocaine Lidocaine (2-(diethylamino)-N -(2,6-dimethylphenyl)acetamide) is a local anesthetic commonly used to relieve pain related to surgical, dental, and gynecological procedures. The therapeutic and toxic effects of lidocaine are directly related to its concentration and metabolites. The toxicity of lidocaine affects primarily the cardiovascular and central nervous systems, and overdoses can result in ventricular

208

CATHODIC TREATMENT OF BDD ELECTRODES AND THEIR USE IN ELECTROANALYSIS

TABLE 8.2 Sildenafil citrate content in Viagra ® pharmaceutical formulations determined by a proposed differential pulse voltammetric (DPV) method using an HT-BDD electrode and by a comparative HPLC method. (Reprinted from Ref. 89.) Determined value (mg) DPV a

HPLC a

Relative errorb (%)

25.7 ± 0.9 51.5 ± 0.5 105.0 ± 0.9

26.1 ± 0.5 52.4 ± 0.8 102.0 ± 0.6

–1.5 –1.7 2.9

Label value (mg) 25 50 100

a Average of 3 measurements. b 100 × [(DPV value—HPLC value)/HPLC value].

arrhythmia. Taking into account the lack of a simple and direct electroanalytical method for lidocaine determination, Oliveira et al. [90] investigated the use of SWV with a BDD electrode for this purpose. According to the authors, a cathodic pretreatment was necessary for conditioning the BDD surface prior to the electroanalytical determinations. Thus, before each analysis the BDD electrode was pretreated in a 0.1 mol L−1 HClO4 solution by applying + 3.2 V versus Ag/AgCl for 30 s (to clean the electrode surface), followed by −2.8 V versus Ag/AgCl, for 30 s. Figure 8.22 shows the SWV responses obtained for different concentrations of lidocaine as well as the corresponding analytical curve. The obtained LOD and LOQ values were 10.0 and 34.4 mg L−1 , respectively. The proposed method was successfully applied in the determination of lidocaine in three different commercial gel formulations, in which lidocaine is mixed with propyleneglycol; the presence of propyleneglycol had no influence on the voltammetric responses, as it can be verified in Table 8.3.

60 Ip (µA)

50

70 60

I (µA)

40 30

50 40 30 20 10 0

20

background curve

(b) 6 8 10 12 2 4 Concentration (10–5 M)

0

10 0

(a) 0.0

0.4

0.8

1.2

1.6

2.0

E (V) Ag/AgCl

Figure 8.22 (a) SWV response obtained for the oxidation of lidocaine at different concentrations in a 0.1 mol L−1 Britton-Robson buffer (pH 2.0); (b) Linear dependence of the oxidation peak current with the lidocaine concentration. SWV conditions: a = 50 mV, Es = 2 mV, f = 150 Hz. BDD electrode area = 0.25 cm2 . (Reprinted from Ref. 90.)

8.5 CONCLUSIONS

209

TABLE 8.3 Lidocaine recovery percentage from samples of commercial pharmaceutical preparations (gels) analyzed using square-wave voltammetry and a cathodically pretreated BDD. (Reprinted from Ref. 90.)

Cream A Cream B Cream C

Label value (mg g−1 )

Found value (mg g−1 )

Recoverya (%)

RSDb (%)

50.0 50.0 50.0

49.8 49.2 48.8

99.6 98.4 97.6

2.1 2.3 2.6

a 100 × (found value/label value); b Relative standard deviation of measurements done in triplicate.

8.4

GOLD DEPOSITION AND STRIPPING

Finally, it should be noted that BDD films have also been used as electrodes to investigate the deposition of different metals, when the electrochemical pretreatment of their surfaces was also found to play a role. Specifically, Holt et al. [91] reported on the deposition of gold on BDD by the reduction of tetrachloroaurate(III). Considering that the adhesion and the electrochemical reactivity of gold particles on diamond likely depended on the physical and chemical nature of the electrode surface, Holt and colleagues investigated the deposition and stripping characteristics of gold metal on a cathodically or anodically pretreated BDD electrode. The cathodic or anodic pretreatment was carried out in dilute aqua regia by applying − 1.2 V (versus SCE) for 10 s or + 2.0 V (versus SCE) for 30 s, respectively. Significant increases in both the gold deposition current and the stripping efficiency were found when the cathodically pretreated BDD electrode was used. The authors reported that the cathodic pretreatment had a dramatic effect on the reduction process of the tetrachloroaurate(III) ion: decrease of the reduction overpotential by about 0.1 V and enhancement of the reduction current by nearly a factor of 2. Additionally and even more significant, the size of the stripping peak increased, enhancing the stripping efficiency. XPS quantitative analyses revealed that the amount of gold deposited on the cathodically pretreated BDD electrode was about 10 times higher than that on an untreated BDD electrode.

8.5

CONCLUSIONS

The cathodic pretreatment of BDD electrodes, which increases the fraction of hydrogen termination on their surfaces, has been shown to enhance the electrochemical activity of the electrodes toward many redox couples and analytes. Consequently, the use of hydrogen-terminated BDD in the electroanalytical determination of several analytes has been found to be quite convenient, since boosted sensitivities and selectivity could be attained. Such effects have been reported in the determination of different pesticides in environmental samples, sweeteners, and total phenols in food samples, and drugs in pharmaceutical samples. Nevertheless, a better characterization of the superficial state of BDD brought on by different types of cathodic pretreatments is still necessary and might contribute to even more significant advancements in the application of this electrode in electroanalysis.

210

CATHODIC TREATMENT OF BDD ELECTRODES AND THEIR USE IN ELECTROANALYSIS

REFERENCES 1. Y.V. Pleskov, A.Y. Sakharova, M.D. Krotova, L.L. Bouilov, B.V. Spitsyn, J. Electroanal. Chem. 1987, 228 , 19–27. 2. A. Fujishima, Y. Einaga, T.N. Rao, D.A. Tryk, Diamond Electrochemistry, Elsevier, Tokyo, 2005. 3. J.C. Angus, Y.V. Pleskov, S.C. Eaton, in Semiconductors and Semimetals: Thin Film Diamond II (C.E. Nebel, J. Ristein, eds.), Elsevier, San Diego, 2004, 77 , pp. 97–120. 4. G.M. Swain, in Semiconductors and Semimetals: Thin Film Diamond II (C.E. Nebel, J. Ristein, eds.), Elsevier, San Diego, 2004, 77 , pp. 121–148. 5. G.M. Swain, in Electroanalytical Chemistry (A.J. Bard, I. Rubinstein, eds.), Marcel Dekker, New York, 2004, 22 , pp. 182–278. 6. O. Chailapakul, W. Siangproh, D.A. Tryk, Sens. Lett . 2006, 4 , 99–119. 7. G. Chen, Talanta 2007, 74 , 326–332. 8. J. Park, V. Quaiserov´a-Mocko, B.A. Patel, M. Novotny, A. Liu, X. Bian, J.J. Galliganb, G.M. Swain, Analyst 2008, 133 , 17–24. 9. K. Peckov´a, J. Musilov´a, J. Barek, Crit. Rev. Anal. Chem. 2009, 39 , 148–172. 10. J.H.T. Luong, K.B. Male, J.D. Glennon, Analyst 2009, 134 , 1965–1979. 11. S. Jolley, M. Koppang, T. Jackson, G.M. Swain, Anal. Chem. 1997, 69 , 4099–4107. 12. M.C. Granger, G.M. Swain, J. Electrochem. Soc. 1999, 146 , 4551–4558. 13. E. Popa, H. Notsu, T. Miwa, D.A. Tryk, A. Fujishima, Electrochem. Solid-State Lett . 1999, 2 , 49–51. 14. H. Notsu, I. Yagi, T. Tatsuma, D.A. Tryk, A. Fujishima, J. Electroanal. Chem. 2000, 492 , 31–37. 15. D.A. Tryk, H. Tachibana, H. Inoue, A. Fujishima, Diam. Rel. Mat . 2007, 16 , 881–887. 16. M. Panizza, I. Duo, P.A. Michaud, G. Cerisola, C. Comninellis, Electrochem. Solid-State Lett . 1999, 3 , 429–430. 17. I. Duo, C. L´evy-Cl´ement, A. Fujishima, C. Comninellis, J. Appl. Electrochem. 2004, 34 , 935–943. 18. H.B. Suffredini, V.A. Pedrosa, L. Codognoto, S.A.S. Machado, R.C. Rocha-Filho, L.A. Avaca, Electrochim. Acta 2004, 49 , 4021–4026. 19. E. Mah´e, D. Devilliers, C. Comminellis, Electrochim. Acta 2005, 50 , 2263–2277. 20. G.R. Salazar-Banda, L.S. Andrade, P.A.P. Nascente, P.S. Pizani, R.C. Rocha-Filho, L.A. Avaca, Electrochim. Acta 2006, 51 , 4612–4619. 21. H. Girard, N. Simon, D. Ballutaud, M. Herlem, A. Etcheberry, Diam. Rel. Mat . 2007, 16 , 316–325. 22. N. Simon, H. Girard, M. Manesse, D. Ballutaud, A. Etcheberry, Diam. Rel. Mat . 2008, 17 , 1371–1375. 23. A. Denisenko, C. Pietzka, A. Romanyuk, H. El-Hajj, E. Kohn, J. Appl. Phys. 2008, 103 , 1–8. 24. M. Wang, N. Simon, C. Decorse-Pascanut, M. Bouttemy, A. Etcheberry, M. Li, R. Boukherroub, S. Szunerits, Electrochim. Acta 2009, 54 , 5818–5824. 25. C.H. Goeting, F. Marken, A. Guti´errez-Sosa, R.G. Compton, J.S. Foord, Diam. Rel. Mat . 2000, 9 , 390–396. 26. F. Marken, R.G. Compton, C.H. Goeting, J.S. Foord, S.D. Bull, S.G. Davies, J. Solid State Electrochem. 2001, 5 , 88–93. 27. I. Yagi, H. Notsu, T. Kondo, D.A. Tryk, A. Fujishima, J. Electroanal. Chem. 1999, 473 , 173–178.

REFERENCES

211

28. S. Ferro, A. De Battisti , Electrochim. Acta 2002, 47 , 1641–1649. 29. K. Hayashi, S. Yamanaka, H. Watanabe, T. Sekiguchi, H. Okushi, K. Kajimura, J. Appl. Phys. 1997, 81 , 744–753. 30. J.B. Cui, J. Ristein, L. Ley, Appl. Phys. Lett . 1998, 81 , 429–432. 31. F. Maier, M. Riedel, B. Mantel, J. Ristein, L. Ley, Phys. Rev. Lett . 2000, 85 , 3472–3475. 32. D. Takeuchi, H. Kato, G.S. Ri, T. Yamada, P.R. Vinod, D. Hwang, C.E. Nebel, H. Okushi, S. Yamasaki, Appl. Phys. Lett . 2005, 86 , 1–3. 33. P.K. Bachmann, W. Eberhardt, B. Kessler, H. Lade, K. Radermacher, D.U. Wiecher, H. Wilson, Diam. Rel. Mat . 1996, 5 , 1378–1383. 34. M. Tachiki, T. Fukuda, K. Sugata, H. Seo, H. Umezawa, H. Kawarada, Jpn. J. Appl. Phys. 2000, 39 , 4631–4632. 35. K.B. Holt, A.J. Bard, Y. Show, G.M. Swain. J. Phys. Chem. B 2004, 108 , 15117–15127. 36. D.A. Tryk, K. Tsunozaki, T.N. Rao, A. Fujishima, Diam. Rel. Mat . 2001, 10 , 1804–1809. 37. M.N. Latto, D.J. Riley, P.W. May, Diam. Rel. Mat . 2000, 9 , 1181–1183. 38. M.N. Latto, G. Pastor-Moreno, D.J. Riley, Electroanalysis 2004, 16 , 434–441. 39. I. Duo, A. Fujishima, C. Comninellis, Electrochem. Comm. 2003, 5 , 695–700. 40. D. Becker, K. J¨uttner, J. Appl. Electrochem. 2003, 33 , 959–967. 41. D. Becker, K. J¨uttner, New Diam. Front. Carbon Technol . 2003, 13 , 67–78. 42. D. Becker, K. J¨uttner, Electrochim. Acta 2003, 49 , 29–39. 43. L. Codognoto, S.A.S. Machado, L.A. Avaca, Diam. Rel. Mat . 2002, 11 , 1670–1675. 44. V.A. Pedrosa, L. Codognoto, L.A. Avaca, Quim. Nova 2003, 26 , 844–849. 45. V.A. Pedrosa, S.A.S. Machado, L.A. Avaca, Anal. Lett . 2006, 39 , 1955–1965. 46. S. Ferro, M. Dal Colle, A. De Battisti, Carbon 2005, 43 , 1191–1203. 47. M. Szameitat, X. Jiang, W. Beyer, Appl. Phys. Lett . 2000, 77 , 1554–1556. 48. J. Ristein, F. Maier, M. Riedel, J.B. Cui, L. Ley, Phys. Stat. Sol. (A) 2000, 181 , 65–76. 49. M.C. Rossi, F. Spaziani, S. Salvatori, G. Conte. Phys. Stat. Sol. A 2003, 199 , 71–76. 50. J. Ristein, M. Riedel, L. Ley, J. Electrochem. Soc. 2004, 151 , E315–E321. 51. J. Ristein, W. Zhang, L. Ley, Phys. Rev. E 2008, 78 , 041602–1–041602-8. 52. J.B. Cui, J. Ristein, L. Ley, Phys. Rev. B 1999, 59 , 5847–5856. 53. S. Kulesza, J. Patyk, F. Rozploch, Chem. Phys. Lett . 2004, 391 , 56–59. 54. M. Hupert, A. Muck, J. Wang, J. Stotter, Z. Cvackova, S. Haymond, Y. Show, G.M. Swain, Diam. Rel. Mat . 2003, 12 , 1940–1949. 55. J. Wang, F. Liu, H. Chen, D. Chen, Mater. Chem. Phys. 2009, 115 , 590–598. 56. G. Pastor-Moreno, D.J. Riley, Electrochem. Comm. 2002, 4 , 218–221. 57. Y.V. Pleskov, Russ. J. Electrochem. 2002, 38 , 1275–1291. 58. S. Wang, G.M. Swain, J. Phys. Chem. C 2007, 108 , 3986–3995. 59. Y.V. Pleskov, J. Anal. Chem., 2000, 55 , 1045–1050. 60. V.A. Pedrosa, H.B. Suffredini, L. Codognoto, S.T. Tanimoto, S.A.S. Machado, L.A. Avaca, Anal. Letters 2005, 38 , 1115–1125. 61. M. Panizza, G. Cerisola, Electrochim. Acta 2005, 51 , 191–199. 62. M.A.Q. Alfaro, S. Ferro, C.A. Martinez-Huitle, Y.M. Vong, J. Braz. Chem. Soc. 2006, 17 , 227–236. 63. C.A. Martinez-Huitle, S. Ferro, Chem. Soc. Rev . 2006, 35 , 1324–1340. 64. L.S. Andrade, L.A.M. Ruotolo, R.C. Rocha-Filho, N. Bocchi, S.R. Biaggio, J. Iniesta, V. Garc´ıa-Garc´ıa, V. Montiel, Chemosphere 2007, 66 , 2035–2043.

212

CATHODIC TREATMENT OF BDD ELECTRODES AND THEIR USE IN ELECTROANALYSIS

65. C.A. Mart´ınez-Huitle, M.A. Quiroz, C. Comninellis, S. Ferro, A. De Battisti, Electrochim. Acta 2004, 50 , 949–956. 66. L.S. Andrade, R.C. Rocha-Filho, N. Bocchi, S.R. Biaggio, J. Iniesta, V. Garc´ıa-Garc´ıa, V. Montiel, J. Hazard. Mater. 2008, 153 , 252–260. 67. A. Kapalka, G. F´oti, C. Comninellis, Electrochim. Acta 2009, 54 , 2018–2023. 68. R.A. Medeiros, A.E. Carvalho, R.C. Rocha-Filho, O. Fatibello-Filho, Anal. Lett . 2007, 40 , 3195–3207. 69. R.A. Medeiros, A.E. Carvalho, R.C. Rocha-Filho, O. Fatibello-Filho, Quim. Nova 2008, 31 , 1405–1409. 70. R.A. Medeiros, A.E. Carvalho, R.C. Rocha-Filho, O. Fatibello-Filho, Talanta 2008, 76 , 685–689. 71. T.A. Ivandini, T.N. Rao, A. Fujishima, Y. Einaga, Anal. Chem. 2006, 78 , 3467–3471. 72. T.A. Ivandini, K. Honda, T.N. Rao, A. Fujishima, Y. Einaga, Talanta 2007, 71 , 648–655. 73. L.S. Andrade, R.C. Rocha-Filho, Q.B. Cass, O. Fatibello-Filho, Electroanalysis 2009, 21 , 1475–1480. 74. D.A. Tryk, T. Kondo, A. Fujishima, in Diamond Electrochemistry (A. Fujishima, Y. Einaga, T.N. Rao, D.A. Tryk, eds.), Tokyo, Elsevier, 2005, pp. 174–217. 75. V. Fisher, D. Gandini, S. Laufer, E. Blank, C. Comninellis, Electrochim. Acta 1998, 44 , 521–524. 76. A. Perret, W. Haenni, N. Skinner, X.M. Tang, D. Gandini, C. Comninellis, B. Correa, G. Foti, Diam. Rel. Mat . 1999, 8 , 820–823. 77. J.M.P. Carrazon, A.D. Recio, L.M.P. Diez, Talanta 1992, 39 , 631–635. 78. L. Codognoto, S.T. Tanimoto, V.A. Pedrosa, H.B. Suffredini, S.A.S. Machado, L.A. Avaca, Electroanalysis 2006, 18 , 253–258. 79. V.A. Pedrosa, L. Codognoto, L. A. Avaca, J. Braz. Chem. Soc. 2003, 14 , 530–535. 80. G.S. Garbellini, G.R. Salazar-Banda, L.A. Avaca, J. Braz. Chem. Soc. 2007, 18 , 1095–1099. 81. D.G. Crosby, Pure Appl. Chem. 1981, 53 , 1051–1080. 82. M.V.S Pinheiro, M.N. Oliveira, A.L.B. Penna, A.Y. Tamine, Int. J. Dairy Technol . 2005, 4 , 193–199. 83. H. Dejmkova, M. Scampicchio, J. Zima, J. Barek, S. Mannino, Electroanalysis 2009, 21 , 1014–1018. 84. T.A. Ivandini, B.V. Sarada, T.N. Rao, A. Fujishima, Analyst 2003, 128 , 924–929. 85. C. Terashima, T.N. Rao, B.V. Sarada, A. Fujishima, Anal. Chem. 2002, 74 , 895–902. 86. C.D. Souza, O.C. Braga, I.C. Vieira, A. Spinelli, Sens. Actuators B 2008, 135 , 66–73. 87. E.R. Sartori, R.A. Medeiros, R.C. Rocha-Filho, O. Fatibello-Filho, J. Braz. Chem. Soc. 2009, 20 , 360–366. 88. B.C. Lourenc¸a˜ o, R.A. Medeiros, R.C. Rocha-Filho, L.H. Mazo, O. Fatibello-Filho, Talanta 2009, 78 , 748–752. 89. E.F. Batista, E.R. Sartori, R.A. Medeiros, R.C. Rocha-Filho, O. Fatibello-Filho, Anal. Letters 2010, 43 , 1046–1054. 90. R.T.S. Oliveira, G.R. Salazar-Banda, V.S. Ferreira, S.C. Oliveira, L.A. Avaca, Electroanalysis 2007, 19 , 1189–1194. 91. K.B. Holt, G. Sabin, R.G. Compton, J.S. Foord, F. Marken, Electroanalysis 2002, 14 , 797–803.

Part IV

Industrial Applications

9 Use of Boron-Doped Diamond Electrode in Electrochemical Generation and Applications of Ferrate Virender K. Sharma, Enric Brillas, Ignasi Sir´es, and Karel Bouzek

9.1

INTRODUCTION

Iron commonly exists in the zero, +2, and +3 oxidation states; however, in certain environments, higher oxidation states of iron such as +4, +5, and +6 can also be obtained [1–4]. Fe(IV) and Fe(V) species have been proposed as reactive intermediates in selective oxygenation of hydrocarbons by iron-induced activation of hydrogen peroxide in organic solvents [5,6]; the ferryl (FeIV = O) and perferryl (FeV = O) species may play an important role in the oxygen activation and transfer reactions mediated by heme and nonheme iron proteins [7–11]. In recent years, iron in the +6 oxidation state, commonly called ferrate (FeVI O4 2− ), has been of great interest because of its role as an oxidant and hydroxylating agent in industrial and water treatment processes, such as the development of a “super iron” battery, the green chemistry synthesis, and the nonchlorine oxidation/disinfection of aqueous effluents for pollutant remediation [12–14]; Fe(VI) provides an environmentally benign, high-energy density battery cathode [15–17], whereas selective oxidations by Fe(VI) can be utilized for synthesizing organic compounds without the release of toxic by-products [18].

Synthetic Diamond Films: Preparation, Electrochemistry, Characterization, and Applications, First Edition. Edited by Enric Brillas and Carlos Alberto Mart´ınez-Huitle.  2011 John Wiley & Sons, Inc. Published 2011 by John Wiley & Sons, Inc.

215

216

USE OF BORON-DOPED DIAMOND ELECTRODE

The aqueous ferrate solutions have a distinctive violet color similar to that of solutions of the permanganate ion. The alkaline solutions of the ferrate ion show a maximum at 510 nm (ε = 1150 ± 25 M−1 cm−1 ) and a shoulder between 275 and 320 nm [19]. The ferrate ion is a powerful oxidizing agent in aqueous media. Under acidic conditions, the reduction potential of the ferrate ion is relatively high compared to that of other oxidants used as disinfectants in water treatment processes, as shown in Table 9.1 [19,20]; however, the reduction potential strongly decreases under more alkaline conditions (Table 9.1). The spontaneous decomposition of ferrate in water leads to the formation of molecular oxygen according to Reaction (9.1) [21]: 2FeO4 2− + 5H2 O → 2Fe(OH)3 + 3/2 O2 + 4OH−

(9.1)

The decomposition rate of ferrate is strongly dependent on the initial ferrate concentration, temperature, pH, and even on the surface properties of the hydrous iron oxide formed upon decomposition. The ferrate ion is an emerging water treatment oxidant, disinfectant, and coagulant, which can address the concerns on disinfection by-products (DBPs) associated with currently used chemicals, such as free chlorine, chloramines, and ozone [22]. Like ozone, Fe(VI) does not react with the bromide ion, and so, the carcinogenic bromate ion is not produced during the treatment of bromide-containing water [22]. Several studies on the reaction of ferrate with a series of pollutants including sulfide, thiourea, cyanides, amines, phenols, and anilines have shown that the destruction of pollutants by ferrate is achieved in seconds to minutes; moreover, it leads to the formation of relatively nontoxic by-products [13]. The reaction rates are pH-dependent; thus, the half-lives determined for the pollutant removal are also pH-dependent [23,24]. Ferrate can also degrade effectively emerging contaminants such as estrogens, bisphenol-A, and sulfonamide antimicrobials present in water [25–30]. Furthermore, the nontoxic byproduct generated from ferrate reaction (i.e., ferric oxide/hydroxide) acts as a powerful TABLE 9.1 Redox potentials for the oxidants/disinfectants used in water treatment [19,20]. Reaction

E ◦ (V/SHE)

OH + H+ + e− ⇔ H2 O OH + e− ⇔ OH− FeO4 2− + 8H+ + 3e− ⇔ Fe3+ + 4H2 O FeO4 2− + 4H2 O + 3e− ⇔ Fe(OH)3 + 5OH− O3 + 2H+ + 2e− ⇔ O2 + H2 O O3 + H2 O + 2e− ⇔ O2 + 2OH− H2 O2 + 2H+ + 2e− ⇔ 2H2 O H2 O2 + 2e− ⇔ 2OH− MnO4 − + 4H+ + 3e− ⇔ MnO2 + 2H2 O MnO4 − + 8H+ + 5e− ⇔ Mn2+ + 4H2 O MnO4 − + 2H2 O + 3e− ⇔ MnO2 + 4OH− HClO + H+ + 2e− ⇔ Cl− + H2 O ClO− + H2 O + 2e− ⇔ Cl− + 2OH− ClO4 − + 8H+ + 8e− ⇔ Cl− + 4H2 O Cl2 + 2e− ⇔ 2Cl− O2 + 4H+ + 4e− ⇔ 2H2 O O2 + 2H2 O + 4e− ⇔ 4OH− ClO2 + e− ⇔ ClO2 −

2.80 1.89 2.20 0.70 2.08 1.24 1.78 0.88 1.68 1.51 0.59 1.48 0.84 1.39 1.36 1.23 0.40 0.95

Oxidant Hydroxyl Radical

•

•

Ferrate (VI) Ozone Hydrogen peroxide Permanganate

Hypochlorite Perchlorate Chlorine Dissolved Oxygen Chlorine Dioxide

9.2 ELECTROCHEMICAL GENERATION OF THE FERRATE ION WITH IRON ANODES

217

coagulant that is suitable for the removal of metals, nutrients, radionuclides, and humic acids [31–33]. Because of the wide range of applications of ferrate, the most economic production ways for individual applications have been sought. Basically, there exist three major approaches for the synthesis of ferrate: (1) dry thermal synthesis, (2) wet chemical synthesis, and (3) electrochemical synthesis. In the thermal synthesis, ferric oxides are heated together with an alkali metal oxide or peroxide to obtain solid ferrate. Different dry syntheses have recently been reviewed [34]. In the wet chemical synthesis, the ferric ion is converted into the ferrate ion by oxidation with hypochlorite in a highly alkaline environment [35]. Ozone and oxone (a mixture of K2 SO4 , KHSO4 , and KHSO5 ) can be used instead of hypochlorite to synthesize ferrate [36,37]. The electrochemical synthetic procedure generally occurs in a concentrated solution of alkali metal hydroxides. Traditionally, two kinds of electrochemical setups have been employed, based on the use of (1) inert anodes (e.g., Pt) able to oxidize soluble Fe3+ ions and (2) iron-based anodes (e.g., grey cast iron, white cast iron, steel, mild steel, etc.) that act as the iron source. More recently, boron-doped diamond (BDD) electrode has been applied to the synthesis of ferrate. The present contribution summarizes different electrochemical synthetic approaches for the generation of the ferrate ion. Several descriptions have been recently reviewed [38]. Hence, this chapter briefly reviews the electrochemical synthetic procedures that are available but mainly focuses on the use of the BDD anode for the electrogeneration of the ferrate ion, as well as on the recent applications of this methodology.

9.2 ELECTROCHEMICAL GENERATION OF THE FERRATE ION WITH IRON ANODES Several studies have been addressed to overcome some of the difficulties associated with the electrochemical synthesis of ferrate [16,17,38–41]. Problems include the formation of a residual passive film on the electrode surface and the extent to which the competitive oxygen evolution reaction (OER) is given at the potential at which ferrate is formed. The current efficiency for the ferrate(VI) generation is very sensitive to the type of electrode pretreatment applied, as well as to the reaction conditions. Therefore, main attention has been paid to the chemical composition, geometry, and mode of activation of the anode, as well as to the electrolyte composition in order to optimize the synthesis of ferrate [38,41]. The influence of other experimental parameters including the temperature, applied current, and electrolysis time has been assessed as well. Various studies have focused on establishing the suitable conditions for electrosynthesizing ferrate when using an iron anode. It has been demonstrated that the presence of high silicon content in grey cast iron allows enhancing the ferrate production, reaching current yields in the range of 20–40% [42]. As can be seen in Figure 9.1, alloys produced by centrifugation led to better results compared to molded alloys, and a content of 2.80% Si produced a current yield of 33%. The degree of porosity of the iron anode also has great influence on the process [43]. For example, Figure 9.2 illustrates the comparative production of ferrate using pellet electrodes. After 1 h of electrolysis, the current yield was much higher using the foil electrode when working at the lowest current density, whereas the pressed iron powder gave the highest yield working over 5.5 mA cm−2 current density [43]. In the latter case, the porous structure of the anode favors the dissolution process, thus accelerating the transformation of the dissolved iron

218

USE OF BORON-DOPED DIAMOND ELECTRODE

40

Current yield (%)

30

20

10

0 2.0

2.4

2.2

2.6

3.2

3.0

2.8

Si content (wt %) Figure 9.1 Dependence of the current yield with the silicon content in grey cast iron for twohour ferrate production using a U-cell with a glass diaphragm. Volume of anodic solution: 50 cm3 ; electrode area: 30 cm2 ; current density: 17–25 mA cm−2 . () Centrifugation, () molded. (Reprinted with permission from Ref. 42.)

100 100 80 Current yield (%)

Current yield (%)

80

60

Pellet

60 40 20

Foil

40

0

0

10 20 30 40 50 60 70 80 j (mA cm–2)

20 Pellet Foil 0

0

100

200 j (mA

300

400

500

cm–2)

Figure 9.2 Current yield as a function of current density after 1 h of electrolysis for (•) iron foil and () pellet electrodes. The hollow circles (◦) correspond to the current yields computed by taking the zero time at the instant at which the electrode potential reached 0.6 V for the pellet electrode. The inset shows the plot for the low current density region. (Reprinted with permission from Ref. 43.)

9.2 ELECTROCHEMICAL GENERATION OF THE FERRATE ION WITH IRON ANODES

219

into ferrate. It is then evident that the maximum current yield varies depending on the physical appearance of the anode used in the electrolysis. Recently, rapid electrochemical preparation of Na2 FeO4 was studied in detail by using an iron wire gauze anode [44]. The ferrate yield (η) was determined upon variation of the NaOH concentration, anodic current density, and electrolysis time. The increased ratio of effective anode surface area to anolyte volume using a thinner anodic chamber enhanced the production rate of ferrate in the anolyte. The dependence of the ferrate concentration, apparent ferrate yield (ηapp ), NaOH concentration, and Fe(III) concentration versus the electrolysis time is presented in Figure 9.3. The curve for the ηapp shows a vertical step because it represents the 1h-average apparent yield. The highest ηapp occurred during the first hour, which corresponded to 0.148 M ferrate. The maximum ferrate concentration (=0.48 M) was obtained at 6 h. But one of the major causes of relatively low ferrate production was the formation of a passive layer of iron oxide on the iron anode surface. Hence, the study was made using anisomeric square pulse with fluctuating frequency to favor the continuous activation of the anode, aiming at obtaining a high ferrate yield [45]. The results in Figure 9.4, which were obtained in a cylindrical electrolytic bath with bipolar membranes, show that the ferrate concentration increased over time in all cases, whereas the corresponding current efficiency always exhibited a progressive decay. This suggests that it was though possible to destroy the passive iron oxide film, but the formation of the passive layer still occurred. Therefore, a low current efficiency and a gradual decrease of the ferrate generation rate were observed [45]. The optimum fluctuating frequency was 2 Hz, as depicted in the inset of Figure 9.4b. Table 9.2 summarizes the current efficiency values reported for the ferrate production under various conditions, using different iron-containing anodes [46–55]. It seems clear that the nature of the anode and the carbon content have a large influence on the current efficiency. Based on the cyclic voltammograms for the formation of ferrate from metallic iron, the following mechanism involving Reactions (9.2) through (9.7) has been proposed 1.0 16

0.8

SFe-[FeO42–]

0.6

[OH–] happ

15

14 0.4

[OH–] (M)

Content(mM) or happ

[FeO42–]

13 0.2 12

0.0 0

1

2

3

4 5 Time (h)

6

7

8

Figure 9.3 Time course of some species and ηapp during a continuous electrolysis using 14-layer iron gauzes for 8 h at 35.0 ± 0.8◦ C. Initial anolyte: 16 M NaOH; volume: 75 cm3 ; estimated total anode surface area: 690 cm2 . The ηapp is the average apparent yield for 1-h intervals. (Reprinted with permission from Ref. 44.)

USE OF BORON-DOPED DIAMOND ELECTRODE

60

30

(a)

25

55

50

20 h (%)

Oxidant(mM FeO42– )

60

(b)

15

2Hz

h (%)

220

50 45

40

0.2Hz

20Hz

0Hz 40 Frequency/Hz

30

10 2Hz 20Hz 0.2Hz 0Hz

5

2Hz 20Hz 0.2Hz 0Hz

20

0 0

1

2

4 3 Time (h)

5

0

6

1

2

3 4 Time (h)

5

6

7

Figure 9.4 (a) Concentration of electrogenerated FeO4 2− and (b) current efficiency with time at 5 mA cm−2 by varying the applied frequency between 0 and 20 Hz for a CS–CMC bipolar membrane electrolysis cell. (Reprinted with permission from Ref. 45.)

[39–41,48,56–60]. Reactions (9.2) and (9.3) represent both the active dissolution of the anode material and the surface layer restructuration, respectively. The reaction of the passive layer of FeOOH with OH− ions results in the breakage of the iron oxide surface and allows the continuous dissolution of the anode material to form FeO2 − by Reaction (9.4). This can be given simultaneously to the oxidation to the FeO3 2− ion from Reaction (9.5a), followed by disproportionation to the ferrate ion from Reaction (9.6). However, the formed FeO2 − ions dissolved in the anolyte may be further anodically oxidized to yield FeO3 2− via Reaction (9.5b). The concerned reaction of oxygen evolution occurs in the vicinity of ferrate formation zone from Reaction (9.7). Fe + 2OH− → Fe(OH)2 + 2e−

(9.2)

−

Fe(OH)2 + OH → FeOOH + H2 O + e

−

FeOOH + OH− → FeO2 − + H2 O

(9.3) (9.4)

FeOOH + 3OH− → FeO3 2− + 2H2 O + e−

(9.5a)

FeO2 − + 2OH− → FeO3 2− + H2 O + e−

(9.5b)

3FeO3 2− + H2 O → 2FeO2 − + FeO4 2− + 2OH−

(9.6)

−

2OH → H2 O + 1/2O2 + 2e

−

(9.7)

This mechanism is generally in good agreement with the oxidation steps undergone by metallic iron to yield Fe(III), but some controversy exists concerning the formation of the ferrate ion from Fe(III). The electrochemical impedance spectroscopy (EIS) study of the system suggests an alternative four-step mechanism; after the oxidation of the zero-valent iron to FeOOH (Reactions 9.2 and 9.3), the reaction can proceed according to Reactions (9.8) and (9.9) [61]: FeOOH + 3OH− → (FeO3 − )ads + 2H2 O + 2e− −

−

−

3(FeO3 )ads + 2OH → FeO2 + 2FeO4

2−

+ H2 O

(9.8) (9.9)

9.2 ELECTROCHEMICAL GENERATION OF THE FERRATE ION WITH IRON ANODES

221

TABLE 9.2 Effect of the operation conditions on the current efficiency for the ferrate production with iron-based anodes. Current efficiency (%)

Operation conditions

Reference

12

Anode: Silver steel with 0.08% C j = 1 mA cm−2 [NaOH] = 9.0 M T = 25◦ C Electrolysis time = 15 min

[46,47]

21

Anode: Grey cast iron j = 2 mA cm−2 [NaOH] = 14 M T = 20◦ C Electrolysis time = 60 min Anode: Mild steel j = 4 mA cm−2 [NaOH] = 14 M T = 20◦ C Electrolysis time = 180 min Anode: Porous magnetite j = 3.3 mA cm−2 [NaOH] = 16 M T = 30◦ C Electrolysis time = 300 minutes Anode: White cast iron j = 5 mA cm−2 [NaOH] = 14 M T = 20◦ C Electrolysis time = 60 min Anode: Pure iron j = 4.4 mA cm−2 [NaOH] = 14 M T = 50◦ C Electrolysis time = 60 min Anode: Grey cast iron j = 4.54 mA cm−2 [NaOH] = 14 M T = 20◦ C Electrolysis time = 60 minutes Anode: Silver steel with 0.9% C j = 1 mA cm−2 [NaOH] = 10 M T = 25◦ C Electrolysis time = 5 min

[48]

42

52.3

64

64

68.5

>70

[49]

[54]

[50]

[51]

[55]

[46,47]

More recently, spectroscopic approaches have also been used to gain information on the processes taking place at the electrode surface, which eventually will help to produce ferrate in a more efficient manner [62]. One of the main drawbacks of the electrochemical synthesis is the instability of ferrate in aqueous medium. Therefore, the molten hydroxides approach, which excludes altogether water from the environment, has been applied to perform the direct anodic oxidation process with an iron anode or with soluble iron species using a platinum anode [63–65]. Under molten hydroxides conditions, the ferrate produced is expected to be in dry stable solid form in the cooled-down reaction mixture. Furthermore, the previously

222

USE OF BORON-DOPED DIAMOND ELECTRODE

proposed mechanism involves a chemical reaction; hence, the increased temperature in the molten state can accelerate the overall generation of ferrate. The electrochemical synthesis of ferrate using this approach was therefore performed at 170, 180, and 200◦ C; the anode passivation was less pronounced during the ferrate(VI) generation in comparison with the experiments in an aqueous environment [63–65], which is important for the continuous production of ferrate. However, the problem related to the competing OER found in aqueous solution was noticed in the molten environment as well. Furthermore, electrolysis operation temperature became the key factor in the process. An excessive temperature increase was detrimental regarding the ferrate yield because of the limited thermal stability of ferrate [66,67]. Conversely, only a few studies have discussed the use of an inert electrode for synthesizing ferrate, although it is advantageous because a bulk Fe3+ solution can be used, which allows avoiding the impact of the anode dissolution kinetics on the overall process. Moreover, only three electrons are needed to generate the ferrate ion, instead of the six electrons required when using zero-valent (i.e., metallic) iron, which reduces the electrical consumption to half the value. The reported synthesis of ferrate using an inert electrode is reviewed hereafter.

9.3 ELECTROCHEMICAL GENERATION OF THE FERRATE ION WITH INERT ANODES In a previous study, cyclic voltammetric (CV) curves were collected using a rotating ringdisk electrode (RRDE), as well as platinum electrodes under stationary conditions [38]. Solutions of 0.03 M Fe(II), Fe(III), and Fe(VI) in a 14 M NaOH electrolyte solution were investigated. CV curves were also recorded in the 14 M NaOH solution using a platinum electrode with the surface covered by cathodically deposited iron. During the potential scan to more positive values, a current plateau in the potential region from +0.060 to +0.675 V versus Hg/HgO was observed in all the studied solutions. This plateau was ascribed to the oxidation of FeO2 − to FeO4 2− . Importantly, the current density of the plateau strongly depended on the electrolyte temperature, whereas it was independent of the rotation speed. This suggests that the mechanism of ferrate formation involves at least one chemical step. Recently, CV curves for the Fe(VI)/Fe(III) system were investigated using a SnO2 -Sb2 O3 /Ti anode in strong basic solution [68,69]. Figure 9.5 shows steady-state CVs obtained with solutions of FeO2 − in 14 M NaOH. They suggest the formation of an intermediate during the oxidation of FeO2 − to FeO4 2− . The slope of the linear relationship in the inset of Figure 9.5 is about 0.028; accordingly, two electrons are transferred in the oxidation step, which is associated with the formation of Fe(V) from Reaction (9.10) and its disproportionation to give the ferrate ion via Reaction (9.11): FeO2 − + 4OH− → FeO4 3− + 2H2 O + 2e− 3FeO4

3−

+ 2H2 O → 2FeO4

2−

−

+ FeO2 + 4OH

(9.10) −

(9.11)

This system was also examined using a powder microelectrode as the anode, which yielded similar conclusions [69]. Not only FeO2 − , but also Fe2 O3 and Fe(OH)3 were used as iron sources, yielding no ferrate ion instead.

9.4 ELECTROCHEMICAL GENERATION OF THE FERRATE ION WITH BORON-DOPED DIAMOND ANODE

223

(a) 0.15 scan rate: 1 m V s–1

j (mA cm–2)

0.10

0.05

0.00

0.0

0.2

0.4

0.6

0.8

j (V) vs Hg/HgO j(V)vs Hg/HgO

(b) 0.12

j (mA cm–2)

0.08

0.65

0.60

0.55 –2

–1

0

1

2

log (il(i1–i))

0.04

scan rate: 0.2 m V s–1 0.00

0.1

0.2

0.3 0.4 0.5 j (V) vs Hg/HgO

0.6

0.7

Figure 9.5 Steady-state cyclic voltammogram at (a) 1 mV s−1 and (b) 0.2 mV s−1 at a SnO2 –Sb2 O3 electrode in 14 M NaOH containing 0.02 M FeO2 − . The arrows indicate the scan direction. The inset gives the linear relationship of ϕ vs. log(i/(il − i)). (Reprinted from Ref. 68.)

9.4 ELECTROCHEMICAL GENERATION OF THE FERRATE ION WITH BORON-DOPED DIAMOND ANODE 9.4.1

Acidic Medium

Original work on the use of the BDD anode for the electrochemical generation of ferrate was carried out in 0.1 M HClO4 [70]. A 0.006 M FeSO4 solution was used to collect the CVs given in Figure 9.6a at different scan rates. The voltammograms exhibited three peaks: two anodic (AI and AII) and one cathodic (CI). The redox couple Fe3+ /Fe2+ is responsible for peaks AI (∼+1.1 V versus Ag/AgCl) and CI (∼+0.4 V versus Ag/AgCl). The peak potential corresponding to AII varied from +2.3 to +2.75 V versus Ag/AgCl

224

USE OF BORON-DOPED DIAMOND ELECTRODE

(a) 4m

CI

0 AI

–4m j (A cm–2)

(a)

(b)

–8m

(c)

–12m (d) –16m

(e)

–20m

(f) (g)

–24m 3.0

AII 2.5

2.0 1.5 1.0 E (V) vs Ag/AgCl

0.5

0.0

(b) CI 0.0 (a) (b)

AI

j (A cm–2)

–1.0m (c) –2.0m

–3.0m AII (d)

–4.0m 2.5

2.0

1.5 1.0 E (V) vs Ag wire

0.5

0.0

Figure 9.6 Cyclic voltammograms at a boron-doped diamond (BDD) electrode. Plot (A): (a) 0.1 M HClO4 alone, or in the presence of 6 mM FeSO4 at scan rates of (b) 10, (c) 50, (d) 100, (e) 250, (f) 500, and (g) 1000 mV s−1 . The electrochemical cell was a single compartment cell with the surface of the BDD electrode exposed at the bottom of the cell through an O-ring supported opening with a Pt mesh counter electrode, and a Ag/AgCl reference electrode (in saturated KCl). Plot (B): (a) acetonitrile with 0.1 M LiClO4 alone, (b) as in (a) but with addition of 6 mM FeSO4 , (c) as in (b) but with addition of 1.13 M of water, and (d) as in (b) but with addition of 2.83 M of water. Scan rate: 100 mV s−1 . A Pt mesh counter electrode and Ag wire pseudo-reference electrode were used. (Reprinted with permission from Ref. 70.)

depending on the scan rate, and it was assigned to the generation of the ferrate ion from the oxidation of Fe3+ in accordance with Reaction (9.12): Fe3+ + 4H2 O → FeO4 2− + 8H+ + 3e−

(9.12)

9.4 ELECTROCHEMICAL GENERATION OF THE FERRATE ION WITH BORON-DOPED DIAMOND ANODE

225

This is supported by the known thermodynamic potential in acidic medium (see Table 9.1) [71]. Interestingly, no cathodic peak was seen for the reduction of Fe(VI) to Fe(III) and the amount of oxygen in the system did not affect the peak. Furthermore, the AII peak was much larger than what expected from Reaction (9.12). This was accounted for by the fact that the generated ferrate ion rapidly decomposed in acidic medium to Fe3+ and oxygen by Reaction (9.13): 2FeO4 2− + 5H2 O → 2Fe3+ + 13/2O2 + 10H+

(9.13)

A similar set of CV experiments was performed by using acetonitrile instead of water as the solvent in order to provide evidence for Reaction (9.12). The CVs of Figure 9.6b show a decrease in the overpotential required by the Fe3+ /Fe2+ redox couple, along with a new anodic peak, AII, at a less positive potential compared to that required for the ferrate formation. This demonstrates the important role of the water content in the system, as deduced from Reaction (9.12). The authors concluded that the initial apparent number of electrons transferred resulted in 3, which further confirmed the occurrence of Reaction (9.12). 9.4.2

Alkaline Medium

The work on the electrochemical generation of ferrate ion using the BDD electrode in alkaline medium has been performed only recently [72,73]. The time course of generated ferrate was monitored in experiments using both, stainless steel and BDD electrodes, as depicted in Figure 9.7. The concentration of ferrate was determined by using the chromite method [35]. When the stainless steel electrode was initially used for 2 h, the concentration of the generated ferrate increased progressively and then the production slowed down to reach a constant value. This can be related to the existence of Reaction (9.14): Fe + 8OH− → FeO4 2− + 4H2 O + 6e−

(9.14)

Oxidant (mM FeO42–)

0.25 0.2 0.15 Anode: SS 0.1

Anode: BDD

0.05 0 0

50

100 150 Time (min)

200

250

Figure 9.7 Variation of the ferrate concentration with time during the electrolysis of hydroxide solutions with () stainless steel and () BDD electrodes. Experimental conditions: 14 M NaOH, at current density of 13 mA cm−2 and 30◦ C. (Reprinted with the permission from Ref. 72.)

226

USE OF BORON-DOPED DIAMOND ELECTRODE

The replacement of the stainless steel anode by a BDD one in the solution treated for 2 h resulted in an increase in the concentration of the ferrate ion; this was possibly due to the oxidation of the iron species accumulated in the solution during the preliminary electro-oxidation process. This finding is interesting because it suggests that BDD can increase the efficiency of the ferrate electrogeneration process. In such studies, the cell voltage did not change, which indicates that the usual phenomenon of passivation and deterioration of the electrode did not occur. The efficiency was found to be dependent on the electrical charge consumed [72]. Experiments were also carried out using the BDD anode, with iron powder and Fe(III) hydroxide as the iron source, at different current densities. Figure 9.8 illustrates that iron powder turned out to be the preferred primary iron material. This suggests that the process to generate ferrate using BDD is determined by the availability of the iron species. Figure 9.8 also shows that an increase in the current density led to a larger production of ferrate. This phenomenon is similar to that observed for the synthesis of other oxidants using BDD (i.e., H2 O2 , S2 O8 2− , etc.) [74–77]. Hydroxyl radicals formed when using BDD electrodes may also be coresponsible for the formation of ferrate in the BDD system [72]. The effect of other parameters, such as the concentration of hydroxide anions and the operation temperature, on the efficiency of ferrate production was assessed as well [73]. The effect of the content of hydroxide anions is shown in Figures 9.9a and 9.9b. The production of ferrate at ≤5 M NaOH was not significant, whereas continuous ferrate yield was obtained as the concentration of hydroxide salt increased up to 14 M NaOH (Figure 9.9a). This may be related to the stability of ferrate after its generation in strong alkaline solution. Figure 9.9b demonstrates that the use of KOH instead of NaOH as the electrolyte markedly increased the amount of ferrate produced. The different stability of the generated ferrate salts, K2 FeO4 and Na2 FeO4 , in the system under alkaline conditions may explain these results. 1.6

Oxidant(mM FeO42−)

1.4 1.2 1 0.8 0.6 0.4 0.2 0 0

500

1000

1500

2000

Current density (A m-2) Figure 9.8 Variation of the ferrate concentration with the current density in the electrolysis with BDD electrodes of hydroxide solutions () saturated with 9.4 mM Fe(OH)3 and () with iron-powder bed. Experimental conditions: 10 M KOH and temperature of 30◦ C. (Reprinted with permission from Ref. 72.)

9.4 ELECTROCHEMICAL GENERATION OF THE FERRATE ION WITH BORON-DOPED DIAMOND ANODE

Oxidant (mM) FeO42−)

0.1

227

(a)

0.08 0.06 0.04 0.02 0 0

Oxidant (mM FeO42−)

0.08

1

2 Q (Ah L–1)

3

4

1

2 Q (Ah L–1)

3

4

(b)

0.06

0.04

0.02

0 0

Figure 9.9 Variation of the ferrate concentration with the electric charged passed during the electrolysis of Fe(OH)3 solutions with BDD electrodes at 13 mA cm−2 at 30◦ C. Plot (a): () 5, () 10, and () 14 M NaOH. Plot (b): () 10 M NaOH and (△) 10 M KOH. (Reprinted with permission from Ref. 73.)

Figures 9.10a and 9.10b show the effect of the electric charge consumed in the electrolysis of iron hydroxide at several temperatures. The former figure presents the initial dependence of the ferrate concentration on the electric charge passed at all temperatures. This fact is very clear at electric charge ≤5 Ah L−1 . Figure 9.10b depicts the change of the ferrate generation yield with the temperature, reaching a maximum at 25◦ C. Two different processes—the solubility of Fe(III) and the stability of ferrate—may be occurring simultaneously to cause such dependence of the ferrate yield on the temperature. At low temperatures, the solubility of Fe(III) is low and a small amount of Fe(III) species is then available to be oxidized, causing a small production of ferrate; conversely, the stability of ferrate decreases with increasing temperature to give low yields of ferrate at too high temperatures (Figure 9.10b). In summary, the current density, electric charged passed (which is related to the electrolysis time), concentration of electrolyte, and temperature have a great influence on the efficiency of the electrosynthesis of ferrate with the BDD anode.

228

USE OF BORON-DOPED DIAMOND ELECTRODE

Oxidant (mM) FeO42−)

0.24

(a)

0.2 0.16 0.12 0.08 0.04 0

0

Oxidant (mM) FeO42−)

0.15

5

10

15 Q (Ah L–1)

20

25

30

(b)

0.1

0.05

0 10

20

30

40

50

60

Temperature (°C) Figure 9.10 Variation of the ferrate concentration (a) with the electric charge passed and (b) with the temperature, during the electrolysis of Fe(OH)3 solutions with BDD electrodes in 10 M KOH at 100 mA cm−2 . Plot (a), temperature: () 17, (△) 23, () 30, (*) 40, and () 57◦ C. Plot (b), ferrate concentrations obtained at an electric charge passed of 1.5 Ah L−1 . (Reprinted with permission from Ref. 73.)

9.5

APPLICATIONS

In contrast to the wide use of chemically produced ferrate, a very limited amount of work has been conducted on the applications of electrochemically generated ferrate. Examples reported in the field of fuel cells as well as in the remediation of water pollutants are described briefly in subsections below on the basis of the different types of electrodes used as the anode. 9.5.1

Common Inert Anodes

In recent years, direct methanol fuel cells (DMFCs) have shown applications in electronic equipments [78]. In alkaline solution, the anodic reaction of the DMFCs is described as follows: CH3 OH + 6OH− → CO2 + 5H2 O + 6e−

(9.15)

9.5 APPLICATIONS

229

Pt or Pt-based alloy anodes are commonly used for methanol oxidation, but these anodes become easily poisoned by the CO-like intermediates. Recent work demonstrated that addition of ferrate ion and Fe(III) in 1 M methanol and 12 M NaOH promoted both the catalytic activity and poison tolerance of Pt [78]. The assessment of the long-term stability of the methanol electro-oxidation in the presence of ferrate and Fe(III) ions under alkaline conditions was carried out through current-time curves at a steady potential. As can be seen in Figure 9.11, there is a fast current decay during the early stages; then, after ∼500 seconds the decay becomes slower in the four electrolytes, and the current progressively tends to zero during the rest of the electro-oxidation of methanol. The initial trends of the curves suggest poisoning of the electrocatalyst. The lowest current was obtained in the absence of ferrate or Fe(III) ions. Interestingly, the oxidation current in the electrolyte containing the decomposed ferrate ion was the largest compared to solutions with a stoichiometric amount of Fe(NO3 )3 and K2 FeO4 . These results were also in agreement with the CVs of the systems, indicating that the reduction product of ferrate ion greatly promoted the electrocatalytic activity of Pt for the oxidation of methanol in alkaline media. 9.5.2

Iron Anodes

The online preparation and use of ferrate in the wastewater treatment at a pilot scale has been only tested recently [79]. Steel was used as the anode in the pilot plant setup for the electrochemical production of ferrate. The steel had an iron content >99% and carbon content in the range of 0.10–0.12%. The thickness of each steel plate was 2 mm and a bipolar configuration of the electrodes was used. The optimum current density was 3.6 mA cm−2 . In the laboratory scale, the possible risk arising from the hydrogen

Normalized current (mA cm–2)

1.0

0.8

0.6

0.4

0.2 d a

0.0 0

b

c 5000

10000 Time (s)

15000

20000

Figure 9.11 Current–time curves for methanol electro-oxidation in 1 M CH3 OH + 12 M NaOH: (a) at −0.041 V versus Hg/HgO, (b) with a Fe(NO3 )3 loading of 203 mg L−1 , (c) with a K2 FeO4 loading of 164 mg L−1 , and (d) with the decomposed K2 FeO4 at −0.080 V versus Hg/HgO. (Reprinted with permission from Ref. 78.)

230

USE OF BORON-DOPED DIAMOND ELECTRODE

production in the electrochemical process was not observed [80]. However, such risks need to be examined for the electrochemical generation of ferrate at the large scale. In the treatment process, different parameters of water quality such as suspended solids (SS), chemical oxygen demand (COD), biological oxygen demand (BOD), and phosphate content were tested. The concentrations in tests were 242–730 mg L−1 for the SS, 523–1125 mg L−1 for the COD, 235–441 mg L−1 for the BOD, and 11.3–18.5 mg L−1 for the phosphate as total P. The results are presented in Figure 9.12. At a maximum ferrate dose of 0.04 mg L−1 , the average removal percentage achieved was 70% for the SS (Figure 9.12a), 40% for the phosphate (Figure 9.12b), 40% for the COD (Figure 9.12c), and 30% for the BOD (Figuer 9.12d). Interestingly, Table 9.3 shows that the performance of a low ferrate dose (0.03 mg Fe L−1 ) could be similar or even better than that obtained with a high amount of ferric sulfate (37 mg Fe L−1 ). The superior performance of ferrate is not surprising, taking into account that it is a much more powerful oxidant compared to the ferric ion. Another significant advantage of ferrate is related to the residual iron concentration in the final effluent. Table 9.3 shows that the use of ferrate leads to only 0.25 mg Fe L−1 in the effluent, whereas the iron concentration in the case of ferric sulfate was much higher (17.2 mg Fe L−1 ). Therefore, the use of a very low dose of ferrate(VI) would generate a much smaller volume and mass of the sludge than that from high doses of ferric salts. This would then be beneficial concerning the reduction of the cost of sludge handling. However, before implementation of the ferrate technology to a full-scale treatment of water and wastewater, the full assessment of the operation cost should be performed. 9.5.3

BDD Anode

Recently, the ferrate ion has been suggested to enhance the electrochemical oxidation of Acid Yellow 36 azo dye (AY36) using the BDD electrode in a system containing Fe(II) ions under acidic conditions [81]. The decolorization of the dye solution using different

100 (a)

80

P removal (%)

SS removal (%)

100

60 40 20 0

60 40 20 0

0

0.01

0.02

0.03

0.04

0.05

0

0.01

0.02

0.03

0.04

0.05

0.01

0.02

0.03

0.04

0.05

100

100 BOD removal (%)

COD removal (%)

(b)

80

(c)

80 60 40 20 0 0

0.01

0.02

0.03

0.04

0.05

(d)

80 60 40 20 0 0

Ferrate dose(mg Fe(VI) L–1)

Figure 9.12 Average removal with ferrate(VI) in the treatment of crude wastewater. (a) Suspended solids, (b) phosphate ion, (c) COD, and (d) BOD. (Reprinted with permission from Ref. 79.)

231

9.5 APPLICATIONS

TABLE 9.3 Comparative performance of crude sewage treatment with ferric sulfate and ferrate(VI)a [79]. Average percentage removal (%)

Residual iron and pH

Chemical and dose

SS

P

COD

BOD

Fe (mg L−1 )

pH

Ferrate(VI) (0.03 mg Fe L−1 ) Ferric sulfate (37 mg Fe L−1 )

79 ± 5

56 ± 1

50 ± 3

30 ± 5

0.25 ± 0.08

9.8 ± 0.1

78 ± 4

59 ± 2

54 ± 5

43 ± 6

17.2

–

a Crude sewage properties: [SS] = 730 mg L−1 , [P] = 18.5 mg L−1 , [COD] = 1125 mg L−1 , [BOD] = 388 mg L−1 , [Fe] = 1.0 mg L−1 , pH = 8.0.

1.0

AY36 (C1/Co)

0.8 (a) 0.6 (b) (c)

0.4

(d) 0.2

(e)

0.0 0

20

40

60

80

100

120

140

160

180

Time (min) Figure 9.13 Effect of the FeSO4 concentration over the decolorisation of 40 mg L−1 of Acid Yellow 36 under an anodic potential of 2.5 V. (a) Electrochemical oxidation process (EOP), and with FeSO4 content of (b) 12, (c) 8, (d) 6, and (e) 1 mM FeSO4 . (Reprinted with permission from Ref. 81.)

amounts of Fe(II) ions is presented in Figure 9.13. The electrochemical oxidation process (EOP) achieved only ∼41% of the AY36 degradation in 180 min. Comparatively, adding Fe(II) ions improved the degradation markedly under the same conditions. The added amount of the Fe(II) ions had a great influence on the degradation of the dye. It is noteworthy that an increase in the concentrations of Fe(II) ions decreased the degradation efficiency. The results obtained in Figure 9.13 were explained by performing CV measurements over the BDD electrode under the same conditions, as illustrated in Figure 9.14. The absence of Fe(II) ions does not yield any redox signal (curve a); similarly, the addition of 12 mM Fe(II) ions does not yield any redox signal (curve b), whereas a signal appeared from Fe(II) contents lower or equal to 8 mM. Three distinctive peaks could be observed: two anodic (AI and AII) and one cathodic (CI) (curve c). The Fe3+ /Fe2+ redox couple was associated with the AI and CI signals, while AII corresponded to the generation of the ferrate ion at ∼+2.3 V from Reaction (9.12). This findings are similar to those reported by Lee et al. (see Section 9.4.1) [70]. The CI peak increased with Fe(II) concentration decreasing from 6 to 1 mM (curves d and e).

232

USE OF BORON-DOPED DIAMOND ELECTRODE

Cl

0.000

(a) (b) (c) (d)

0.002

j (A cm–2)

0.004

AI

0.006 0.008

(e) AII

0.010 0.012 0.014 0.016 3.0

2.5

2.0

1.5 1.0 E (V) vs Ag/AgCl

0.5

0.0

Figure 9.14 Cyclic voltammograms at BDD in 0.1 M HClO4 for ferrate ion generation in the following conditions: (a) without FeSO4 , and with FeSO4 loading of: (b) 12, (c) 8, (d) 6, and (e) 1 mM. Scan rate: 100 mV s−1 . (Reprinted with permission from Ref. 81.)

H2O BDD Oxidation products

a) Oxidation products

H++e– Fe2+

Pollutants

e) BDD c)

b) BDD(•OH) f) Pollutants

[Fe(VI)] d)

Fe2+

Fe3+

Fe3+

Fe3+

Figure 9.15 Proposed mechanism for the electrochemical oxidation of organic compounds with the simultaneous ferrate ion formation in the system. (Reprinted with permission from Ref. 81.)

Results of Figures 9.13 and 9.14 clearly demonstrate the role of Fe(II) ions in the enhancement of the electrochemical oxidation of AY36 using a BDD electrode. The mechanism given in Figure 9.15 suggests that the ferrate ion is the species responsible for such an enhancement. In the proposed mechanism, step (a) produces hydroxyl radical, which oxidizes the AY36 (step (b)). The generation of the ferrate ion occurs in step (c), so that AY36 can also be oxidized by this species (step (d)), which shows the enhancement effect of the Fe(II) ions. The Fe(II) ions are regenerated by the reduction of the Fe(III) ions in step (e). A certain amount of Fe(II) is oxidized by hydroxyl radicals (step (f)). It seems that steps (b) and (d) give rise to a synergistic effect for degrading AY36 at low concentrations of the Fe(II) ions by avoiding minimizing steps (e) and (f).

REFERENCES

9.6

233

CONCLUSIONS

Significant advances in terms of the optimization of parameters such as the anode material and operating conditions have been made regarding the electrochemical synthesis of ferrate. However, the mechanism of the Fe(III) oxidation to yield ferrate still needs to be fully understood. The use of classical electrochemical techniques may not be sufficient to explore the mechanism because the reaction of Fe(III) to Fe(VI) is overlapped by oxygen evolution. Hence, spectroscopic approaches are currently being used to gain insight on the processes occurring at the electrode surface. Two relatively new synthesis approaches—the employment of molten hydroxides as the electrolysis environment and the utilization of inert anodes—are currently being pursued to minimize the influence of the anode material and the oxygen evolution in synthesizing ferrate. Studies reporting the use of BDD anode to synthesize ferrate are forthcoming. The efficiency of the synthesis is greatly influenced by the availability of oxidizable iron species in the reaction solution. The potential role of ferrate in the electrochemical oxidation of methanol and dyes using inert anodes have been suggested, but mechanistic understanding of the system is still missing. The in situ electrochemical ferrate production method is promising for treating pollutants in water and wastewater treatment facilities. 9.7

ACKNOWLEGMENTS

V.K. Sharma would like to acknowledge the support of United States National Science Foundation (CHE 0706834). V.K. Sharma and K. Bouzek also acknowledge the partial support of NATO collaborative linkage grant (CBP.EAP.CLG.983119). K. Bouzek acknowledges the support of the Ministry of Education, Youth and Sports of the Czech Republic (project No. ME 890).

REFERENCES 1. J.D. Rush, B.H.J. Bielski, J. Am. Chem. Soc. 1986, 108 , 523–525. 2. C. Jeannot, B. Malaman, R. Gerardin, B. Oulladiaf, J. Solid State Chem. 2002, 165 , 266–277. 3. J.F. Berry, E. Bill, E. Bothe, S.D. George, B. Mienert, F. Neese, K. Wieghardt, Science 2006, 312 , 1937–1941. 4. F.T. De Oliveira, A. Chanda, D. Banerjee, X. Shan, S. Mondal, L. Que, Jr., E.L. Bominaar, E. M¨unck, T.J. Collins, Science 2007, 315 , 835–838. 5. H.C. Tung, C. Kang, D.T. Sawyer, J. Am. Chem. Soc. 1992, 114 , 3445–3455. 6. J.T. Groves, J. Inorg. Biochem. 2006, 100 , 434–447. 7. J. Kaizer, E.J. Klinker, N.Y. Oh, J.-U. Rohde, W.J. Song, A. Stubna, J. Kim, E. M¨unck, W. Nam, L. Que, Jr., J. Am. Chem. Soc. 2004, 126 , 472–473. 8. E.Y. Tshuva, S.J. Lippard, Chem. Rev . 2004, 104 , 987–1011. 9. S.V. Kryatov, E.V. Rybak-Akimova, S. Schindler, Chem. Rev . 2005, 105 , 2175–2226. 10. N.Y. Oh, Y. Suh, M.J. Park, M.S. Seo, J. Kim, W. Nam, Angew. Chem. Int. Ed . 2005, 44 , 4235–4239. 11. X. Shan, L. Que, Jr., J. Inorg. Biochem. 2006, 100 , 421–433. 12. J.D. Carr, P.B. Kelter, A. Tabatabai, D. Splichal, J. Erickson, C.W. McLaughlin, in Proceedings of the Fifth Conference on Water Chlorination: Chemistry, Environmental Impact and Health

234

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.

USE OF BORON-DOPED DIAMOND ELECTRODE

Effects (R.L. Jolley, R.J. Bull, W.P. Davis, S. Katz, M.H. Roberts, V.A. Jacobs, eds.), Lewis Publishers, Chelsea, MI, 1985, Vol. 5, pp. 1285–1298. V.K. Sharma, Adv. Environ. Res. 2002, 6 , 143–156. V.K. Sharma, F. Kazama, J. Hu, A.K. Ray, J. Water Health 2005, 3 , 45–58. S. Licht, B. Wang, S. Ghosh, Science 1999, 285 , 1039–1042. S. Licht, X. Yu, in Ferrates: Synthesis, Properties, and Applications in Water and Wastewater Treatment (V.K. Sharma, ed.), Vol. 985, ACS Symposium Series, 2008, pp. 2–51. S. Licht, X. Yu, in Ferrates: Synthesis, Properties, and Applications in Water and Wastewater Treatment (V.K. Sharma, ed.), Vol. 985, ACS Symposium Series, 2008, pp. 197–256. L. Delaude, P. Laszlo, J. Org. Chem. 1996, 61 , 6360–6370. V.K. Sharma, in Advances in Control of Disinfection By-Products in Drinking Water Systems (A. Nikolau, L. Rizzo, H. Selcuk, eds.), Nova Science Publishers, New York, 2007, pp. 370–380. E. Brillas, I. Sir´es, M.A. Oturan, Chem. Rev . (doi. 9.1021/cr900136g). H. Goff, R.K. Murmann, J. Am. Chem. Soc. 1971, 93 , 6058–6065. V.K. Sharma, S.K. Mishra, N. Nesnas, Environ. Sci. Technol. 2006, 40 , 7222–7227. V.K. Sharma, Water Sci. Technol . 2004, 49 , 69–74. R.A. Yngard, V.K. Sharma, J. Filip, R. Zboˇril, Environ. Sci. Technol. 2008, 42 , 3005–309. J.Q. Jiang, J. Hazard. Mater. 2007, 146 , 617–623. V.K. Sharma, Chemosphere 2008, 73 , 1379–1386. V.K. Sharma, X.Z. Li, N. Graham, R.-A. Doong, J. Water Supply: Res. Technol.- AQUA 2008, 57 , 419–426. G.A.K. Anquandah, V.K. Sharma, J. Environ. Sci. Health, Part A: Toxic/Hazard. Subst. Environ. Eng. 2009, 44 , 62–66. V.K. Sharma, G.A.K. Anquandah, N. Nesnas, Environ. Chem. Lett . 2009, 7 , 115–119. V.K. Sharma, G.A.K. Anquandah, R.A. Yngard, H. Kim, J. Fekete, K. Bouzek, A.K. Ray, D. Golovko, J. Environ. Sci. Health, Part A: Toxic/Hazard. Subst. Environ. Eng. 2009, 44 , 423–442. M.E. Potts, D.R. Churchwell, Water Environ. Res. 1994, 66 , 107–109. U.M. Joshi, R. Balasubramanian, V.K. Sharma, in Ferrates: Synthesis, Properties, and Applications in Water and Wastewater Treatment (V.K. Sharma, ed.), Vol. 985, ACS Symposium Series, 2008, pp. 466–476. Y. Lee, S.G. Zimmermann, A.T. Kieu, U. Von Gunten, Environ. Sci. Technol. 2009, 43 , 3831–3838. Y.D. Perfiliev, V.K. Sharma, in Ferrates: Synthesis, Properties, and Applications in Water and Wastewater Treatment (V.K. Sharma, ed.), Vol. 985, ACS Symposium Series, 2008, pp. 112–123. G.W. Thompson, L.T. Ockerman, J.M. Schreyer, J. Am. Chem. Soc. 1951, 73 , 1379–1381. M.D. Johnson, US patent No. 5746994, 1998. Y.D. Perfiliev, E.M. Benko, D.A. Pankratov, V.K. Sharma, S.K. Dedushenko, Inorg. Chim. Acta 2007, 360 , 2789–2791. Z. M´acov´a, K. Bouzek, J. H´ıveˇs, V.K. Sharma, R.J. Terryn, J.C. Baum, Electrochim. Acta 2009, 54 , 2673–2683. K. Bouzek, M.J. Schmidt, A.A. Wragg, Collect. Czech. Chem. Commun. 2000, 65 , 133–140. F. Lapicque, G. Valentine, Electrochem. Commun. 2002, 4 , 764–766. X. Yu, S. Licht, J. Appl. Electrochem. 2008, 38 , 731–742.

REFERENCES

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.

235

V. Lescuras-Darrou, F. Lapicque, G. Valentin, J. Appl. Electrochem. 2002, 32 , 57–63. M. De Koninck, D. B´elanger, Electrochim. Acta 2003, 48 , 1435–1442. W. He, J. Wang, C. Yang, J. Zhang, Electrochim. Acta 2006, 51 , 1967–1973. Y. Ren, Z. Chen, Y. Geng, R. Chen, X. Zheng, Chem. Eng. Process. 2008, 47 , 708–715. A. Denver, D. Plecher, J. Appl. Electrochem. 1996, 26 , 815–821. A. Denver, D. Plecher, J. Appl. Electrochem. 1996, 26 , 823–827. K. Bouzek, I. Rouˇsar, J. Appl. Electrochem. 1996, 26 , 919–923. K. Bouzek, I. Rouˇsar, J. Appl. Electrochem. 1993, 23 , 1317–1322. K. Bouzek, I. Rouˇsar, M.A. Taylor. J. Appl. Electrochem. 1996, 26 , 925–931. K. Bouzek, I. Rouˇsar. J. Appl. Electrochem. 1997, 27 , 679–684. K. Bouzek, M.J. Schmidt, A.A. Wragg, Electrochem. Commun. 1999, 1 , 370–374. K. Bouzek, M.J. Schmidt, A.A. Wragg, J. Chem. Technol. Biotechnol . 1999, 74 , 1188–1194. Z. Ding, C. Yang, Q. Wu, Electrochim. Acta 2004, 49 , 3155–3159. J. Xu, J. Wang, W. Yang, C. Chen, J. Zhang, Dianhuaxue 2004, 10 , 87–93. M. Alsheyab, J.-Q. Jiang, C. Stanford, J. Environ. Manage. 2009, 90 , 1350–1356. F. Beck, R. Kaus, M. Oberst, Electrochim. Acta 1985, 30 , 173–183. J.-Y. Zou, D.-T. Chin, Electrochim. Acta 1988, 33 , 477–485. K. Bouzek, I. Rouˇsar, H. Bergmann, K. Hertwig, J. Electroanal. Chem. 1997, 425 , 125–137. K. Bouzek, M. Nejezchleba, Collect. Czech. Chem. Commun. 1999, 64 , 2044–2060. H.B. Shao, J.M. Wang, W.C. He, J.Q. Zhang, C.N. Cao, Electrochem. Commun. 2005, 7 , 1429–1433. Z. M´acov´a, K. Bouzek, V.K. Sharma, in Ferrates: Synthesis, Properties, and Applications in Water and Wastewater Treatment (V.K. Sharma, ed.), Vol. 985, ACS Symposium Series, 2008, pp. 52–67. J. H´ıveˇs, M. Benov´a, K. Bouzek, V.K. Sharma, Electrochem. Commun. 2006, 8 , 1737–1740. J. H´ıveˇs, Z. M´acov´a, M. Benov´a, K. Bouzek, J. Electrochem. Soc. 2008, 155 , E113–E119. J. H´ıveˇs, M. Benov´a, K. Bouzek, J. Sitek, V.K. Sharma, Electrochem. Acta 2008, 54 , 203–208. J. Madar´asz, R. Zboˇril, Z. Homonnay, V.K. Sharma, G. Pokol, J. Solid State Chem. 2006, 179 , 1426–1433. L. Machala, R. Zboˇril, V.K. Sharma, J. Filip, O. Schneeweiss, Z. Homonnay, J. Phys. Chem. B 2007, 111 , 4280–4286. C.Z. Zhang, Z. Liu, F. Wu, L.-J. Lin, F. Qi, Electrochem. Commun. 2004, 6 , 1104–1109. C.Z. Zhang, H.B. Deng, T. Zhao, F. Wu, W. Liu, S. Cai, K. Yang, V.K. Sharma, in Ferrates: Synthesis, Properties, and Applications in Water and Wastewater Treatment (V.K. Sharma, ed.), Vol. 985, ACS Symposium Series, 2008, pp. 81–93. J. Lee, D.A. Tryk, A. Fujishima, S.-M. Park, Chem. Commun. 2002, 5 , 486–487. R.H. Wood, J. Am. Chem. Soc. 1958, 80 , 2038–2041. P. Ca˜nizares, M. Arc´ıs, C. S´aez, M.A. Rodrigo, Electrochem. Commun. 2007, 9 , 2286–2290. C. S´aez, M.A. Rodrigo, P. Ca˜nizares, AIChE J . 2008, 54 , 1600–1607. K. Serrano, P.A. Michaud, C. Comninellis, A. Savall, Electrochim. Acta 2002, 48 , 431–436. M.S. Saha, T. Furuta, Y. Nishiki, Electrochem. Commun. 2004, 6 , 201–204. P. Ca˜nizares, F. Larrondo, J. Lobato, M.A. Rodrigo, C. S´aez, J. Electrochem. Soc. 2005, 152 , D191–D196.

236

USE OF BORON-DOPED DIAMOND ELECTRODE

77. P. Ca˜nizares, C. S´aez, A. S´anchez-Carretero, M.A. Rodrigo, J. Appl. Electrochem. 2009, 39 , 2143–2149. 78. Z. Xu, J. Hu, Z. Yan, S. Yang, J. Zhou, W. Lu, Electrochim. Acta 2009, 54 , 3548–3552. 79. J.-Q. Jiang, C. Stanford, M. Alsheyab, Sep. Purif. Technol. 2009, 68 , 227–231. 80. M. Alsheyab, J.-Q. Jiang, C. Stanford, J. Chem. Health Safety 2008, 15 , 16–20. 81. M. Villanueva-Rodr´ıguez, A. Hern´andez-Ram´ırez, J.M. Peralta-Hern´andez, E.R. Bandala, M.A. Quiroz-Alfaro, J. Hazard. Mater. 2009, 167 , 1226–1230.

10 Electrochemical Oxidation of Organic Compounds Induced by Electro-Generated Free Hydroxyl Radicals on BDD Electrodes Agnieszka Kapałka, Helmut Baltruschat, and Christos Comninellis

10.1

INTRODUCTION

In the last decade, the advanced oxidation processes (AOPs) have been demonstrated as an effective means of wastewater treatment [1]. The main principle of AOPs is based on oxidation of organic pollutants using hydroxyl radicals that are highly reactive and destroy organics with a very high rate and low selectivity. AOPs are classified depending on the reaction phase or HO • generation method, generally classified as chemical (Fenton processes), electrochemical, sonochemical (ultrasound processes), and photochemical methods (UV/H2 O2 , UV/O3 photolysis, photo-Fenton, TiO2 photocatalysis) [1]. The electrochemical method for the oxidation of target organic pollutants is a technology for treatment of dilute wastewater (COD < 5 gl−1 ). In this method, organic pollutants are oxidized by intermediacy of hydroxyl radicals electro-generated from water on high Synthetic Diamond Films: Preparation, Electrochemistry, Characterization, and Applications, First Edition. Edited by Enric Brillas and Carlos Alberto Mart´ınez-Huitle. © 2011 John Wiley & Sons, Inc. Published 2011 by John Wiley & Sons, Inc.

237

238

ELECTROCHEMICAL OXIDATION OF ORGANIC COMPOUNDS

oxidation power anodes. An ideal anode for electrochemical oxidation of organic compounds is a boron-doped diamond electrode (BDD). The aim of this work is to elucidate the principles of electrochemical oxidation of organic compounds induced by electro-generated free hydroxyl radicals on BDD electrode. The following points will be treated: — Influence of anode material on the reactivity of electrolytic hydroxyl radicals — Electro-generation and detection of quasi-free hydroxyl radicals on BDD anode — Concentration profile of quasi-free hydroxyl radicals during the oxygen evolution reaction and electro-oxidation of organic compounds on BDD — Kinetic model of organics mineralization on BDD — Electrochemically induced mineralization of organics by molecular oxygen on BDD 10.2 INFLUENCE OF ANODE MATERIAL ON THE REACTIVITY OF ELECTROLYTIC HYDROXYL RADICALS In the electrochemical mineralization (EM) reactions oxygen is transferred from water to the organic pollutant using electrical energy. According to the generally accepted mechanism of the EM in acid media [2], water is first discharged at the anode active sites M, producing hydroxyl species (either adsorbed species, if the potential is low or, if the potential is high enough, even free radicals), as seen in Equation 10.1: M + H2 O → M(HO • ) + H+ + e−

(10.1)

These electrogenerated hydroxyl species are involved in the mineralization of organic pollutants R (see Equation 10.2): R + M(HO • ) → M + mineralization products + H+ + e−

(10.2)

This reaction (Equation (10.2) is in competition with the side reaction of the anodic discharge of hydroxyl species to dioxygen (see Equation 10.3): 1 M(HO • ) → M + O2 + H+ + e− 2

(10.3)

The activity of these electrolytic hydroxyl species is strongly linked to their interaction with the electrode surface M [3]. As a general rule, the weaker the interaction, the lower the electrochemical activity is toward oxygen evolution (high O2 overvoltage anodes) and the higher the chemical reactivity is toward organics oxidation. Based on this approach, we can classify the different anode materials according to their oxidation power in acid media as it is shown in Table 10.1 [4]. This table shows that the attainable oxidation potential of the anode is directly related to the overpotential for oxygen evolution and to the adsorption enthalpy of hydroxyl species on the anode surface (i.e., for a given anode material, the higher the O2 overvoltage, the higher its oxidation power is). A low oxidation power anode is characterized by a strong electrode-hydroxyl species interaction resulting in a high electrochemical activity for the oxygen evolution reaction

10.2 INFLUENCE OF ANODE MATERIAL

TABLE 10.1

239

Oxidation power of the anode material in acid media. (Adapted from Ref. 4.)

Electrode RuO2 − TiO2 (DSA-Cl2 ) IrO2 − Ta2 O5 (DSA-O2 ) Ti/PtOx Ti/PbO2 Ti/SnO2 -Sb2 O5 p–Si/BDD

Oxidation potential of organics (V)∗

Overpotential of O2 evolution (V)∗∗

1.4–1.7 1.5–1.8 1.7–1.9 1.8–2.0 1.9–2.2 2.2–2.6

0.18 0.25 0.3 0.5 0.7 1.3

Oxidation power of the anode

∗ Depending on the applied current densities ∗∗ Considering an overpotential for O evolution of 0.1 mA cm−2 2

(low overvoltage anode) and to a low chemical reactivity for organics oxidation (low current efficiency for organics oxidation). A typical low oxidation power anode is the IrO2 -based electrode. Concerning this anode, it has been demonstrated using differential electrochemical mass spectrometry (DEMS) [5] that the interaction between IrO2 and hydroxyl species is so strong that a higher oxidation state oxide is formed as an intermediate. This higher oxide can act as mediator for both organics oxidation and oxygen evolution. On low oxidation power anodes, oxidation of organics occurs with a low current efficiency but with a high selectivity. Typically, in electrosynthesis, a high selectivity is achieved under potentiostatic conditions. However, on low oxidation power anodes, a high selectivity can be achieved also under galvanostatic conditions [6]. This can be explained by the fact that on these anodes, under galvanostatic conditions, the potential is “buffered” by the oxygen evolution reaction, which is the competing side reaction. Under these so-called pseudo-potentiostatic conditions, the working potential is fixed by the nature (oxidation power) of electrode material, which is related to the oxygen evolution overpotential (i.e., the higher the oxygen evolution overpotential, the higher the oxidation power of the anode (Table 10.1)). The principle of the anode potential “buffering” induced by the oxygen evolution reaction is shown in Figure 10.1 This figure shows that during 11 IrO2

10 9 j (mA cm–2)

8 7 6 5

Dj

4 3 2 1 0 0.0

DE 0.5

1.0 1.5 E (V) vs SHE

2.0

Figure 10.1 Principle of buffering the anode potential induced by the oxygen evolution reaction using steady-state current potential curve recorded on Ti/IrO2 in 1 M HClO4 + 100 mM i-propanol. (Adapted from Ref. 6.)

240

ELECTROCHEMICAL OXIDATION OF ORGANIC COMPOUNDS

oxidation of isopropanol on IrO2 , a large increase in current results in a small change of anode potential due to potential buffering by the side reaction of oxygen evolution. In fact, a change of one decade in current induces a shift of only 40 mV of potential. On the contrary to low oxidation power anode, the high oxidation power anode is characterized by a week electrode-hydroxyl species interaction resulting in a low electrochemical activity for the oxygen evolution reaction (high overvoltage anode) and to a high chemical reactivity for organics oxidation (high current efficiency for organics oxidation). Boron-doped diamond based anode is a typical high oxidation power anode. Here, the electrode-hydroxyl interaction is so weak that hydroxide radicals are formed. The evidence for the formation of hydroxyl radicals on BDD was found by means of spin trapping, hydroxylation of salicylic acid, and formation of hydrogen peroxide [7], as discussed next.

10.3 ELECTRO-GENERATION AND DETECTION OF QUASI-FREE HYDROXYL RADICALS ON BDD ELECTRODE 10.3.1

Hydroxyl Radicals Spin Trapping

The principle of the spin trapping method is to produce a stable adduct by allowing a specific scavenger to react with a less stable radical. As shown in Equation 10.4, the spin trap 5.5-dimethyl-1-pyrroline-N-oxide (DMPO) reacts with hydroxyl radicals to produce a stable adduct, detectable by electron spin resonance (ESR). H3C H3C

+

HO•

+

N

O− DMPO spin trap

H3C H3C

N

OH

O• DMPO hydroxyl radical spin adduct

(10.4)

The main advantage of using DMPO is that it exhibits different ESR spectra with hydroxyl radical and singlet oxygen. The rate constant between DMPO and hydroxyl radicals is equal to 4.3 × 109 mol−1 s−1 [8]. Figure 10.2 shows ESR spectrum obtained after 2 hours of electrolysis of 10 mM DMPO in 1 M HClO4 at 0.1 mA cm−2 on BDD. This spectrum reveals the existence of the DMPO hydroxyl radical spin adduct (see Equation 10.4). The ESR hyperfine couplings of the DMPO adduct produced by electrolysis (obtained by simulation) are aN = aH = 14.95 G. These hyperfine coupling constants are typical of the spin adduct DMPO-OH [9,10], indicating that hydroxyl radicals are indeed produced during electrolysis on BDD electrode.

10.3.2

Trapping by Salicylic Acid

Salicylic acid (SA) is usually used for trapping free hydroxyl radicals in biological system. In fact, SA reacts with hydroxyl radicals to form dihydroxylated aromatic compounds [7]. The main products of the hydroxylation of salicylic acid with hydroxyl radicals are

10.3 ELECTRO-GENERATION AND DETECTION OF QUASI-FREE HYDROXYL RADICALS

241

Figure 10.2 Electron spin resonance of DMPO adduct obtained after electrolysis of 8.8 mM DMPO solution in 1 M HClO4 for 2 h on BDD electrode at 0.1 mA cm−2 . (Adapted from Ref. 7.)

2,3- and 2,5-dihydroxybenzoic acids (DHBA) (see Equation 10.5). COOH

COOH

OH

COOH

OH

HO•

OH +

OH

HO

2,3-Dihydroxybenzoic acid

2,5-Dihydroxybenzoic acid

(10.5)

Figure 10.3 shows the concentration profile of salicylic acid and its conversion during electrolysis on BDD anode. As expected, the hydroxylation of salicylic acid results in formation of dihydroxylated intermediates. High 2,5- and 2,3-DHBA formation yields cannot be achieved because of its further oxidation to catechol and aliphatic acids. On BDD, the current density has almost no influence on the salicylic conversion as well as

100

6

4

60

X (%)

Conc. (mol m–3)

80

40 2 20 0

0 0

0.4

0.8

1.2

1.6

2

Q (Ah L–1) Figure 10.3 Oxidation of salicylic acid in 1 M HClO4 on BDD anode at 20 A m−2 ; (•) salicylic acid, () 2,5-DHBA, () 2,3-DHBA, () conversion of salicylic acid. (Reprinted from Ref. 7.)

242

ELECTROCHEMICAL OXIDATION OF ORGANIC COMPOUNDS

the DHBA selectivity indicating that the chemical reaction of salicylic acid hydroxylation takes place. 10.3.3

Competitive Reactions

When a mixture of organic compounds is present in solution, one can consider that there is a competition for hydroxyl radicals between the organics. Figure 10.4 shows the results of electrolysis performed on BDD anode in an equimolar solution of formic acid and oxalic acid. It can be seen that the concentration of formic acid decreases much faster than the concentration of oxalic acid. In fact, only when the concentration of formic acid is very low, oxalic acid starts to be oxidized with a significant rate. These results indicate that formic acid competes with oxalic acid for hydroxyl radicals. Indeed, the rate constant between hydroxyl radicals and formic acid (k = 108 M−1 s−1 ) is higher by two orders of magnitude than the rate constant between hydroxyl radicals and oxalic acid (k = 106 M−1 s−1 ) [11]. Such a difference in k between these acids explains why the oxidation of formic acid occurs first on BDD. 10.3.4

Formation of Hydrogen Peroxide

If the interaction between the electrode surface and hydroxyl radicals is weak, hydrogen peroxide is expected to be formed by combination of two hydroxyl radicals, as seen in Equation 10.6: HO • + HO • → H2 O2

(10.6)

Figure 10.5 shows generation of hydrogen peroxide during electrolysis of 1 M HClO4 on BDD anode at the current densities from 23 to 160 mA cm−2 . Depending on the applied current density, the concentration of hydrogen peroxide, after 1 h, reaches a plateau ranging from 0.3 to 1 mM. The obtained steady state H2 O2 concentration (plateau) is

Figure 10.4 Evolution of concentration of formic and oxalic acids during electrolysis of a mixture of 0.5 M formic acid and 0.5 M oxalic acid in 1 M HClO4 on BDD at 238 A m−2 . (Reprinted from Ref. 7.)

10.3 ELECTRO-GENERATION AND DETECTION OF QUASI-FREE HYDROXYL RADICALS

243

1.0

[H2O2] (mol m–3)

0.8

0.6

0.4

0.2

0.0 0

5

10

15

Time (h) Figure 10.5 Production of H2 O2 at different current densities; (♦) 23 mA cm−2 , () 47 mA cm−2 , (△) 95 mA cm−2 , (x) 160 mA cm−2 during electrolysis of 1 M HClO4 on BDD electrode. (Reprinted from Ref. 7.)

due to the further anodic oxidation of hydrogen peroxide to oxygen. The initial current efficiency of H2 O2 formation reaches a value up to 0.5% [12]. Therefore, these results show that upon anodic polarization, a significant amount of hydroxyl radicals is formed at BDD electrode. Considering that free hydroxyl radicals are formed in acid aqueous solution according to Reaction (10.7), it is enlightening to calculate the thermodynamic standard potential of its formation using Equation 10.8 [13]. • − + H+ H2 O(l) ⇄ HO(aq) (aq) + e

E◦ = −

Gr◦ zF

(10.7) (10.8)

where E ◦ (V) is the standard thermodynamic potential of reaction, Gr ◦ (kJ mol−1 ) is the standard Gibbs free energy of the reaction, and z is the number of transferred electrons (z = 1). Taking the standard Gibbs free energy of liquid water formation as −237.178 kJ mol−1 and that of hydroxyl radicals in the aqueous state as −7.74 kJ mol−1 [13], one obtains the thermodynamic standard potential for HO • formation equal to 2.38 V.1 On BDD electrodes, the onset potential of oxygen evolution reaction is about 2.3 V (see Figure 10.6). The fact that O2 is evolved close to the thermodynamic potential of HO • formation (and at potentials positive of it) may indicate that electro-generated hydroxyl radicals on BDD surface are quasi-free. This assumption is justifiable, taking into account the inert nature of diamond surface, which contains closely packed sp3 carbon atoms, and a lack of adsorption sites [14]. As shown in Figure 10.6, the oxygen evolution reaction on BDD occurs with a high overpotential with respect to 1 A somewhat higher value of 2.73 V has also been suggested in literature. [P. Wardman, J. Phys. Chem. Ref. Data, 1989, 18, 1637-1755] For a concentration of 10 µM, which is the steady state concentration obtained only of at high current densities (see below), the equilibrium concentration would be 2.4 V, similar to the value given above.

244

ELECTROCHEMICAL OXIDATION OF ORGANIC COMPOUNDS

BDD IrO2

8

j (mA cm–2)

7 6 5 4

•OH/H

Ir(V)/Ir(IV)

3 2

2O

O2/H2O

1 0 0.0

1.0

2.0

3.0

E (V) vs. SHE

Figure 10.6 Oxygen evolution reaction on IrO2 and BDD electrode in acid media. ◦ = 1.23 V versus SHE), but is very thermodynamic potential for O2 formation (EOER ◦ • = 2.38 V versus SHE). close to the thermodynamic potential of HO formation (EHO Therefore, it might be concluded that formation of this active intermediate (H2 O/HO • ) determines the overpotential for OER on BDD. Similarly, on IrO2 , OER proceeds close to the thermodynamic potential of Ir(V)/Ir(IV) redox couple (Figure 10.6), as it was demonstrated by DEMS measurements [5].

10.4 CONCENTRATION PROFILE OF HYDROXYL RADICALS ON BDD ELECTRODE It is interesting to determine the concentration profile of hydroxyl radicals during oxygen evolution and organics electro-oxidation on BDD electrode. To do that, one should consider that anodically polarized BDD electrode generates quasi-free hydroxyl radicals, which mediate the oxidation processes in the vicinity of the electrode surface. 10.4.1

HO • Concentration Profile during Oxygen Evolution

In the absence of organic compounds, electro-generated free hydroxyl radicals can react with each other to form hydrogen peroxide, as given by Equation 10.6. Hydrogen peroxide can be further oxidized to oxygen either by its direct discharge on the electrode surface (Equation 10.9) or via assistance of hydroxyl radicals (Equation 10.10): H2 O2 → O2 + 2H+ + 2e− •

H2 O2 + 2HO → O2 + 2H2 O

(10.9) (10.10)

Assuming that the first step (Equation 10.6) is the rate determining step and that four hydroxyl radicals (four electrons) are needed for oxygen evolution, by applying onedimensional Fick law referred to the molecular diffusion in a stagnant layer, the mass balance for hydroxyl radicals in an element of width x can be expressed as [15]:     dcHO • dcHO • 2 • • • − 4kHO cHO • x = −DHO (10.11) −DHO dx x dx x +x

10.4 CONCENTRATION PROFILE OF HYDROXYL RADICALS ON BDD ELECTRODE

245

where DHO • (m2 s−1 ) is the diffusion coefficient of HO • , kHO • (m3 mol−1 s−1 ) is the rate constant of Reaction 10.6, and cHO • (mol m−3 ) is the concentration of HO • . If x goes to 0, we can write: DHO •

d2 cHO • 2 = 4kHO • cHO • dx 2

(10.12)

The boundary conditions are obtained by assuming that far from the electrode (x = ∞) concentration of hydroxyl radicals is 0, whereas at the electrode (x = 0) there is a surface s concentration of hydroxyl radicals (cHO • ): cHO • = 0 at x = ∞

(10.13)

s at x = 0 cHO • = cHO •

(10.14)

Considering these boundary conditions, the solution of differential Equation 10.12 gives the concentration profile of hydroxyl radicals, as a function of the distance x from the electrode surface, during oxygen evolution. cHO • =

2kHO •



3DHO • 2  • x + 2k3DHO • cs • HO

(10.15)

HO

The gradient of the concentration can be expressed as: dcHO • = dx

kHO

•



−3DHO • 3  • x + 2k3DHO • cs • HO

(10.16)

HO

Using Equation 10.16, it is possible to calculate the flux of hydroxyl radicals at the electrode surface, as given by Equation 10.17: [JHO ]x =0 = −DHO •

•



dcHO • dx



x =0

 s 3 = 1.63 (cHO • ) kHO • DHO •

(10.17)

The flux of hydroxyl radicals at the electrode surface can also be expressed in terms of the current density j : [JHO • ]x =0 =

j F

(10.18)

Thus, comparing Equation 10.17 with Equation 10.18, the surface concentration of hydroxyl radicals during oxygen evolution (OER) can be obtained (see Equation 10.19):

s cHO •

OER

=

 3

j2 2.67F 2 kHO • DHO •

(10.19)

246

ELECTROCHEMICAL OXIDATION OF ORGANIC COMPOUNDS

Figure 10.7 Simulated concentration profile of hydroxyl radicals during oxygen evolution according to Equations 10.15 and 10.19; j = 300 A m−2 ; DHO = 2.2 × 10−9 m2 s−1 ; kHO = 5.5 × 109 M−1 s−1 [10]. (Reprinted from Ref. 15.)

Figure 10.7 shows the concentration profile of hydroxyl radicals during oxygen evolution as a function of the distance from the electrode surface. It can be seen that for current density of j = 300 A m−2 , the reaction layer thickness is about one micrometer, whereas the maximum (surface) concentration of hydroxyl radicals reaches the value of several tenths of µM.

10.4.2 HO • Concentration Profile during Electro-Oxidation of Organic Compound By analogy to Equations 10.11 through 10.19, it is possible to determine the concentration profile of hydroxyl radicals during oxidation of organic compounds under these hypotheses: •

Oxidation of organic compounds R proceeds only via assistance of hydroxyl radicals in the vicinity of the electrode surface. R + zHO • → oxidation products

(10.20)

Concentration of organic compound is high enough to be considered as a constant in the reaction layer. • Oxygen evolution, via H2 O2 oxidation, is negligible. •

The two latter hypotheses apply to the charge-transfer controlled region.

10.4 CONCENTRATION PROFILE OF HYDROXYL RADICALS ON BDD ELECTRODE

247

Under these hypotheses, the mass balance for hydroxyl radicals leads to Equations 10.21 and 10.22:     dcHO • dcHO • • • • − zkR cHO cR x = −DHO (10.21) −DHO dx x dx x +x DHO •

d2 cHO • = zkR cHO • cR dx 2

(10.22)

where z is the number of electrons involved in oxidation of organic compound R, kR (m3 mol−1 s−1 ) is the rate constant of Reaction 10.20, and cR (mol m−3 ) is the concentration of organic compound. Thus, solving Equation 10.22 with the boundary conditions given in Equations 10.13 and 10.14, one can obtain the concentration profile of hydroxyl radicals during oxidation of organic compounds (Equation 10.23):

cHO • =

s cHO •

 

zkR cR exp − x DHO •



(10.23)

The gradient of the concentration can be expressed as: dcHO • s = −cHO • dx



  zkR cR zkR cR exp − x DHO • DHO •

(10.24)

Using Equation 10.24, the flux of hydroxyl radicals at the electrode surface is given by Equation 10.25:

s zkR cR DHO • [JHO • ]x =0 = cHO •

(10.25)

Comparing Equation 10.25 with Equation 10.18, the surface concentration of hydroxyl radicals during oxidation of organic compounds R is obtained (Equation 10.26): s cHO • = R

j √ F zkR cR DHO •

(10.26)

In the presence of organic compound, concentration of hydroxyl radicals decreases exponentially (Equation 10.23) with the distance from the electrode surface. The thickness of the reaction layer (reaction cage) depends on the root of the concentration of organic compound, the rate constant of organics oxidation (via hydroxyl radicals), and the applied current density. As a typical example, Figure 10.8 shows the simulated HO • concentration profile during oxidation of formic acid (0.25–1 M) at 300A m−2 . It can be seen that the higher is the formic acid concentration, the lower is the surface concentration of HO • , and the smaller is the thickness of the reaction layer. For investigated concentrations of formic acid, the thickness of the reaction layer (reaction cage) drops down to barely tenths of nanometers, which is significantly lower as compared with that in the absence of organics (Figure 10.7). Figure 10.9 shows the comparison of HO • concentration profile for 1 M formic acid, methanol, and ethanol. Depending on the rate constant, the thickness of the reaction layer varies between few nanometers and tenths of nanometers.

248

ELECTROCHEMICAL OXIDATION OF ORGANIC COMPOUNDS

Figure 10.8 Simulated concentration profile of hydroxyl radicals during oxidation of (1) 1 M, (2) 0.75 M, (3) 0.5 M, and (4) 0.25 M formic acid, according to Equations 10.23 and 10.26; j = 300 A m−2 ; DHO = 2.2 × 10−9 m2 s−1 ; kHCOOH = 1.3 × 108 M−1 s−1 [10]; z = 2. (Reprinted from Ref. 15.)

Figure 10.9 Simulated concentration profile of hydroxyl radicals during oxidation of 1 M (1) formic acid, (2) methanol, and (3) ethanol according to Equations 10.23 and 10.26; j = 300 A m−2 ; DHO = 2.2 × 10−9 m2 s−1 ; kHCOOH = 1.3 × 108 M−1 s−1 ; kCH3OH = 9.7 × 108 M−1 s−1 ; kC2H5OH = 1.9 × 109 M−1 s−1 [10]; z = 2, 6, and 12 for formic acid, methanol, and ethanol, respectively. (Reprinted from Ref. 15.)

10.5

KINETIC MODEL OF ORGANICS OXIDATION ON BDD ANODE

Under electrolysis regime, electrogenerated hydroxyl radicals (Equation 10.1) are the intermediates for both the main reaction of organics oxidation (Equation 10.2) and the side reaction of oxygen evolution (Equation 10.3). Considering this simplified reaction scheme (Equations 10.1, 10.2, and 10.3), a kinetic model is proposed based on following

249

10.5 KINETIC MODEL OF ORGANICS OXIDATION ON BDD ANODE

suppositions: (1) adsorption of the organic compounds at the electrode surface is negligible; (2) all organics have the same diffusion coefficient D; and (3) the global rate of the electrochemical mineralization of organics is a fast reaction and it is controlled by mass transport of organics to the anode surface. The consequence of this last assumption is that the rate of the mineralization reaction is independent of the chemical nature of the organic compound present in the electrolyte. Under these conditions, the limiting current density for the electrochemical mineralization of an organic compound (or a mixture of organics) under given hydrodynamic conditions can be written as Equation 10.27 [3]: jlim (t) = 4Fkm COD(t)

(10.27)

where jlim is the limiting current density for organics mineralization (A m−2 ), F is the Faraday constant (C mol−1 ), km is the mass transport coefficient (m s−1 ) and COD is the chemical oxygen demand (mol O2 m−3 ). At the beginning of electrolysis, at time t = 0, the initial limiting current density (jlim ) is given by: 0 jlim = 4Fkm COD0

(10.28)

where COD0 is the initial chemical oxygen demand. Working under galvanostatic conditions, two different operating regimes are defined: at japplied < jlim the electrolysis is controlled by the applied current, whereas at japplied > jlim it is controlled by the mass transport control. 10.5.1

Electrolysis under Current Limited Control ( japplied < jlim )

In this operating regime, the current efficiency is 100% and the rate of COD removal is constant and can be written as: r =α

0 jlim 4F

(10.29)

where α is the dimensionless current density defined as: α=

japplied with 0 < α < 1 0 jlim

(10.30)

Using Relation 10.28, the rate of COD removal (Equation 10.30) can be given by: r = αkm COD0

(10.31)

It is necessary to consider the mass-balances over the electrochemical cell and the reservoir to describe the temporal evolution of COD in the batch recirculation reactor system given in Figure 10.10. Considering that the volume of the electrochemical reactor VE (m3 ) is much smaller than the reservoir volume VR (m3 ), we can obtain from the mass-balance on COD for the electrochemical cell the following relation: QCODout = QCODin − αkm ACOD0

(10.32)

250

ELECTROCHEMICAL OXIDATION OF ORGANIC COMPOUNDS

CO2 O2 W1

W2 F2

F1

R3

R4

P3

P4 _

+

R1

E

R2

R5

R6 P1

P2

Figure 10.10 Scheme of the two-compartment electrochemical flow cell; R = reservoirs; P = pumps; E = electrochemical cell with membrane; W = heat exchangers, F = gas flow controllers. (Reprinted from Ref. 4.)

where Q is the flow-rate (m3 s−1 ) through the electrochemical cell, CODin and CODout are the chemical oxygen demands (mol O2 m−3 ) at the inlet and at the outlet of the electrochemical cell, respectively, and A is the anode area (m2 ). For the well-mixed reservoir (Figure 10.10), the mass balance on COD can be expressed as: Q(CODout − CODin ) = VR

dCODin dt

(10.33)

Combining Equations 10.32 and 10.33 and replacing CODin by the temporal evolution of chemical oxygen demand COD, we obtain: COD0 Akm dCODin = −α dt VR

(10.34)

Integrating this equation subject to the initial condition COD = COD0 at t = 0 gives the evolution of COD(t) with time in this operating regime (japplied < jlim ):  Akm COD(t) = COD 1 − α t VR 0

(10.35)

This behavior persists until a critical time (tcr ), at which the applied current density is equal to the limiting current density, which corresponds to: CODcr = αCOD0

(10.36)

10.5 KINETIC MODEL OF ORGANICS OXIDATION ON BDD ANODE

251

Substituting Equation 10.36 in Equation 10.35 it is possible to calculate the critical time (Equation 10.37): tcr = 10.5.2

1 − α VR α Akm

(10.37)

Electrolysis under Mass Transport Control ( japplied > jlim )

In this operating regime, the current efficiency is 100% and the rate of COD removal is constant and can be written as: When the applied current exceeds the limiting one (japplied > jlim ), secondary reactions (such as oxygen evolution) start to proceed, resulting in a decrease of the instantaneous current efficiency (see Equation 10.38): ICE =

COD(t) jlim = japplied αCOD0

(10.38)

This regime is realized either at: 0 1. japplied < jlim when the electrolysis is continued over the critical time (COD < 0 αCOD at t > tcr ), or 0 2. japplied > jlim when applied current exceeds its initial limiting value (COD < αCOD0 at any finite time)

In these cases, the COD mass balances on the anodic compartment of the electrochemical cell E and the reservoir R2 (Figure 10.10) can be expressed as: dCOD Akm COD =− dt VR

(10.39)

Integration of this equation from t = tcr to t, and COD = αCOD0 to COD(t) leads to:  Akm 1−α at t > tcr t+ 1. COD(t) = αCOD0 exp − VR α  Akm 0 t 2. COD(t) = COD exp − VR

(10.40) (10.41)

From Equations 10.38, 10.40 and 10.41 the instantaneous current efficiency, ICE, is now given by:  Akm 1−α at t > tcr 1. ICE = exp − t+ VR α  Akm 1 2. ICE = exp − t α VR

(10.42) (10.43)

A graphical representation of the proposed kinetic model is given in Figure 10.11. In order to verify it, an anodic oxidation of various organic compounds has been performed on BDD. Figure 10.12 shows both the experimental and predicted values (continuous

252

ELECTROCHEMICAL OXIDATION OF ORGANIC COMPOUNDS

(a)

COD0

A

COD(t ) = COD0(1–

B αAkm t) VR

COD(t) = αCOD0exp ( – CODcr =

i appl

Akm 1–α ) VR t + α

4Fk m

0 (b) ICE(%)

t

100

ICE = exp ( –

o (1–α) Qcr = i cr km

Ak m 1–α ) t+ VR α

0 t

Figure 10.11 Evolution of (a) COD and (b) ICE in function of time (or specific charge); (A) represents the charge transport control; (B) represents the mass transport control. (Reprinted from Ref. 4.)

1.0

2500 ICE/–

0.8

COD (ppm)

2000

0.6 0.4 0.2

1500

0.0 0

5

10

15

Specific charge (Ah L–1)

1000

500

0 0

5

10

15

Specific charge (Ah L–1)

Figure 10.12 Evolution of COD and ICE (inset) in function of specific charge for different organic compounds: (x) acetic acid, () isopropanol, (o) phenol, (△) 4-chlorophenol, (♦) 2-naphtol; i = 238 A m−2 ; T = 25◦ C; Electrolyte: 1 M H2 SO4 ; the solid line represents model prediction. (Reprinted from Ref. 4.)

line) of both ICE and COD evolution with the specific electrical charge passed during the anodic oxidation of different classes of organic compounds (acetic acid, isopropanol, phenol, 4-chlorophenol, 2-naphtol). This figure demonstrates that the electrochemical treatment is independent on the chemical nature of the organic compound. Furthermore, there is an excellent agreement between the experimental data and predicted values from proposed model. Figure 10.13 shows the influence of current density on both ICE

10.6 ELECTROCHEMICALLY INDUCED MINERALIZATION OF ORGANIC COMPOUNDS

253

1.0

2000 ICE /–

0.8

COD (ppm)

1500

0.6 0.4 0.2 0.0

1000

0

5

10

15

Specific charge (Ah L–1) 500

0 0

5

10 Specific charge (Ah

15

L–1)

Figure 10.13 Influence of the applied current density, (x) 119 A m−2 , (o) 238 A m−2 , (♦) 476 A m−2 on the evolution of COD and ICE (inset) during electrolysis of 5 mM 2-naphtol in 1 M H2 SO4 on BDD; T = 25◦ C; the solid line represents model prediction. (Reprinted from Ref. 4.)

and COD evolution with the specific electrical charge passed during the galvanostatic oxidation of a 5 mM 2-naphtol in 1 M H2 SO4 at different current densities (119–476 A m−2 ). As previously, an excellent agreement between the experimental and predicted values is observed.

10.6 ELECTROCHEMICALLY INDUCED MINERALIZATION OF ORGANIC COMPOUNDS BY MOLECULAR OXYGEN Electrochemical mineralization of organic compounds on BDD anode not only involves hydroxyl radicals but also molecular oxygen present in air/oxygen-saturated aqueous organic solutions [16]. The direct evidence for this process was found using differential electrochemical mass spectrometry (DEMS). Figure 10.14 shows a DEMS measurement, in which the cyclic voltammogram (CV) and the mass spectrometric cyclic voltammograms (MSCV) for C16 O2 , 18 O2 , 16 O18 O, C18 O2 , and C16 O18 O were simultaneously recorded during oxidation of acetic acid solution (1) deaerated and (2) saturated with isotopically labeled 18 O2 . Figures 10.14e and 10.14f show that isotopically labeled C18 O2 and C16 O18 O are evolved in parallel to 18 O2 diminution (Figure 10.14c) indicating that oxygen dissolved in the solution is involved in mineralization of acetic acid. The rate of C18 O16 O is higher than that of C18 O2 due to the higher probability of C18 O16 O formation, because one atom of oxygen (16 O) comes from water. It is interesting to note that ionic current of C18 O2 reaches a limitation at higher potentials (Figure 10.14e). The limiting current may indicate that oxidation of acetic acid by 18 O2 (similarly, 16 O18 O formation in Figure 10.14d) is limited by diffusion of 18 O2 to the electrode surface. Under investigated conditions, 7% of additional, non-Faradaic CO2 was evolved. This non-Faradaic enhancement of the acetic acid electro-oxidation proceeds most likely through the sequence of reactions that are initiated by HO • formed on the electrode

254

ELECTROCHEMICAL OXIDATION OF ORGANIC COMPOUNDS

CV

I f (mA)

4

1–2

2

(a) 0 MSCV, m/z = 44 (C16O2)

I i (nA)

6

1–2

3 (b) 0 1.0

2

16

I i (nA)

MSCV, m/z = 36 (18O2)

I j (pA) 0.8

12 1 (c) 8

0.6 75

MSCV, m/z = 34 (16O18O)

I i (pA)

2 50

25 27

(d)

1

MSCV, m/z = 48 (C18O2)

I i (pA)

2 18

9 (e)

1

0

I i (nA)

0.4

MSCV, m/z = 46 (C16O18O) 2

0.2 1

(f) 0.0 1.8

2.0

2.2 2.4 E (V) vs RHE

2.6

2.8

Figure 10.14 Simultaneously recorded cyclic voltammogram (CV) (a) and mass spectrometric CV (MSCV) for (b) C16 O2 (m/z = 44), (c) 18 O2 (m/z = 36), (d) C16 O18 O (m/z = 34), (e) C18 O2 (m/z = 48), and (f) C16 O18 O (m/z = 46) of (1) deaerated and (2) 18 O2 -saturated solution of 50 mM acetic acid; scan rate 10 mV s−1 , flow rate 5 μls−1 , electrolyte 1 M HClO4 , T = 25◦ C. (Adapted from Ref. 16.)

10.6 ELECTROCHEMICALLY INDUCED MINERALIZATION OF ORGANIC COMPOUNDS

255

surface. Such a reaction scheme was proposed in radiolysis of hydrocarbons (RH). It has been reported that during the ionizing irradiation of aqueous solution of organic compounds (upon x-ray or γ -ray), molecular oxygen enhances the oxidation processes [17–19]. This occurs via addition of molecular oxygen to an organic free radical (R • ) resulting in formation of an organic peroxy radical (RO2 • ), which can participate in subsequent reactions. The organic free radical (R • ) is formed via dehydrogenation of hydrocarbons (RH) initiated by HO • formed during radiolysis of water. Therefore, by analogy between radiation-induced oxidation of organic compounds and our system, in which HO • are generated electrochemically, a general model of electrochemically induced oxidation of organic compounds via molecular oxygen on BDD electrode can be proposed, as shown in Figure 10.15. Such a non-Faradaic enhancement of electrolysis opens the possibilities for designing a less energy consuming electrochemical mineralization process in which current is applied mainly to initiate formation of HO • , whereas complete degradation of pollutant proceeds via reaction with molecular oxygen in aerated solutions (Figure 10.15, reactions 1–4). To do so, it is necessary to ensure an efficient transport of molecular O2 to the interface and thus the organic radicals (via porous electrodes). The key factors of the efficient mineralization process are the stability of organic radicals (reaction 5 in Figure 10.15, corresponding to formation of 18 O16 O in Figure 10.14d) and its reactivity toward molecular oxygen dissolved in solution (reaction 3 in Figure 10.15). It can be further concluded that such electrochemically induced activation process may occur also on other electrode materials on which organic radicals are formed during a direct discharge of organic compounds at the electrode surface.

RO•2 (4) (5)

(3) O2 (aq)

O2 R•

RO

1/2O2 (2)

RH (6)

H2O

(1)

H+

HO• e–

H+

e–

BDD electrode surface Figure 10.15 Simplified diagram for electrochemically induced oxidation of organic compounds via molecular oxygen dissolved in aerated solution on boron-doped diamond electrode; (1) water discharge to hydroxyl radical HO • ; (2) dehydrogenation of organic compound RH via HO • and formation of free organic radical R • ; (3) addition of molecular oxygen to R • resulting in formation of an organic peroxy radical RO2 • ; (4) decomposition of RO2 • leading to regeneration of HO • and formation of RO; (5) decomposition of RO2 • to R • ; (6) oxygen evolution, a side reaction. (Adapted from Ref. 16.)

256

10.7

ELECTROCHEMICAL OXIDATION OF ORGANIC COMPOUNDS

CONCLUSIONS

On boron-doped diamond electrodes, organic compounds are oxidized by intermediacy of quasi-free hydroxyl radicals electro-generated from water at high anodic potentials. These hydroxyl radicals initiate chain reactions in which molecular oxygen dissolved in aerated aqueous solution contributes to mineralization of organic compounds at ambient temperature. In general, oxidation of organic compounds on BDD proceeds with very high current efficiency, which makes BDD an ideal anode material for electrochemical mineralization of organic wastes. 10.8

EXERCISES

1. Calculate the concentration of hydroxide radicals (DHO • = 2.2 × 10−9 m2 s−1 ) at the surface of BDD electrode (x = 0), at a distance of 0.1 µm and at 0.5 cm from the electrode surface during oxygen evolution in 1 M HClO4 at 200 A m−2 , T = 25◦ C. Consider that the rate determining step of the process is the following reaction: HO • + HO • → H2 O2

kHO • = 5.5 × 106 m3 mol−1 s−1

2. Calculate the concentration of hydroxide radicals (DHO • = 2.2 × 10−9 m2 s−1 ) at the surface of the BDD electrode (x = 0) during oxygen evolution under galvanostatic conditions in 1 M HClO4 at 25◦ C at current densities of: a) 300 A m−2 b) 100 A m−2 c) 10 A m−2 Consider that the rate determining step of the process is this reaction: HO • + HO • → H2 O2

kHO • = 5.5 × 106 m3 mol−1 s−1

3. The anodic oxidation of 1 M HCOOH solution in 2 M HClO4 has been carried out at 25◦ C under galvanostatic conditions (300 A m−2 ) using BDD electrodes. Calculate the concentration of hydroxide radicals (DHO • = 2.2 × 10−9 m2 s−1 ) at the surface of the BDD electrode (x = 0) after: a) Conversion of 30% of HCOOH to CO2 b) Conversion of 60% of HCOOH to CO2 c) Conversion of 90% of HCOOH to CO2 Consider that the process is controlled by the reactions: HO • + HO • → H2 O2

kHO • = 5.5 × 106 m3 mol−1 s−1

HCOOH + 2HO • → CO2 + 2H2 O

kHCOOH = 1.3 × 105 m3 mol−1 s−1

4. Calculate the initial concentration of hydroxide radicals (DHO • = 2.2 × 10−9 m2 s−1 ) at the surface of the BDD electrode (x = 0) during the anodic oxidation of the following organic compounds under galvanostatic conditions (100 A m−2 ) in 2 M HClO4 at T = 25◦ C:

10.8 EXERCISES

257

a) 1 M HCOOH (kHCOOH = 1.3 × 105 m3 mol−1 s−1 ) b) 1 M CH3 OH (kCH3OH = 9.7 × 105 m3 mol−1 s−1 ) c) 1 M C2 H5 OH (kC2H5OH = 1.9 × 106 m3 mol−1 s−1 ) 5. The mass transfer coefficient of an electrolytic cell, under given hydrodynamic conditions, has been estimated to be 1.2 × 10−5 m s−1 using the ferri-ferrocyanide (FFC) redox couple. Calculate: a) The mass transfer coefficient of formic acid, oxidized at the same conditions b) The limiting current density for the oxidation of 0.1 M formic acid aqueous solution given that DFFC = 1.1 × 10−9 m2 s−1 and DHCOOH = 1.7 × 10−9 m2 s−1 . 10.8.1

Solutions

• 1. .a) HO concentration at the surface of the electrode (x = 0):  j2 3 s cHO • = 2 OER 2.67F kHO • DHO •  2002 3 = = 5.1 × 10−2 mol m−3 2.67 · 964852 · 5.5 · 106 · 2.2 · 10−9

b) HO • concentration at a distance of 0.1 µm from the electrode (x = 0.1 µm = 10−7 m): cHO • =

3DHO • 2   • 2kHO • x + 2k3DHO s •c • HO

HO

3 · 2.2 · 10−9 −2 −3 =  2 = 1.4 × 10 mol m  3·2.2·10−9 2 · 5.5 · 106 10−7 + 2·5.5·10 6 ·5.1·10−2 c) HO • concentration at a distance of 0.5 cm from the electrode (x = 0.5 cm = 5 × 10−3 m): cHO • =

3 · 2.2 · 10−9 −11 mol m−3  2 = 2.4 × 10  −9 3·2.2·10 2 · 5.5 · 106 5 · 10−3 + 2·5.5·10 6 ·5.1·10−2

• 2. a . ) HO concentration at the surface of the electrode (x = 0) for galvanostatic conditions at 300 A m−2 :  j2 3 s cHO • = 2 OER 2.67F kHO • DHO •  3002 3 = = 6.7 × 10−2 mol m−3 2.67 · 964852 · 5.5 · 106 · 2.2 · 10−9

258

ELECTROCHEMICAL OXIDATION OF ORGANIC COMPOUNDS

b) HO • concentration at the surface of the electrode (x = 0) for galvanostatic conditions at 100 A m−2 :

s cHO •

OER

=

 3

2.67 ·

964852

1002 = 3.2 × 10−2 mol m−3 · 5.5 · 106 · 2.2 · 10−9

c) HO • concentration at the surface of the electrode (x = 0) for galvanostatic conditions at 10 A m−2 :

s cHO •

OER

=

 3

2.67 ·

964852

102 = 6.9 × 10−3 mol m−3 · 5.5 · 106 · 2.2 · 10−9

3. .a) During oxidation of an organic compound: s cHO • = R

j √ F zkR cR DHO •

with

cR = cR0 (1 − X )

where X is the conversion rate; cR0 = 1M = 103 mol m−3 HCOOH + 2OH • → CO2 + 2H2 O z = 2 b) HO • concentration at the surface of the electrode (x = 0) after conversion of 30% of HCOOH to CO2 : s cHO • = R

300 j

= √ 5 F zkR cR DHO • 96485 2 · 1.3 · 10 · 103 (1 − 0.3) · 2.2 · 10−9

= 4.9 × 10−3 mol m−3

c) HO • concentration at the surface of the electrode (x = 0) after conversion of 60% of HCOOH to CO2 : s cHO • = R

300

= 6.5 × 10−3 mol m−3 5 96485 2 · 1.3 · 10 · 103 (1 − 0.6) · 2.2 · 10−9

d) HO • concentration at the surface of the electrode (x = 0) after conversion of 90% of HCOOH to CO2 : s cHO • = R

300 = 1.3 × 10−2 mol m−3

5 96485 2 · 1.3 · 10 · 103 (1 − 0.9) · 2.2 · 10−9

10.8 EXERCISES

259

4. .a) HO • concentration at the surface of the electrode (x = 0) during anodic oxidation of 1 M (103 mol m−3 ) HCOOH (kR = 1.3 × 105 m3 mol−1 s−1 ) at 100 A m−2 : HCOOH + 2OH • → CO2 + 2H2 O j s cHO • = √ R F zkR cR DHO • =

z =2

100 √ 96485 2 · 1.3 · 105 · 103 · 2.2 · 10−9

= 1.4 × 10−3 mol m−3

b) HO • concentration at the surface of the electrode (x = 0) during anodic oxidation of 1 M (103 mol m−3 ) CH3 OH (kR = 9.7 × 105 m3 mol−1 s−1 ) at 100 A m−2 : CH3 OH + 6OH • → CO2 + 5H2 O z =6 100 s = 2.9 × 10−4 mol m−3 cHO • = √ R 96485 6 · 9.7 · 105 · 103 · 2.2 · 10−9 c) HO • concentration at the surface of the electrode (x = 0) during anodic oxidation of 1 M (103 mol m−3 ) C2 H5 OH (kR = 1.9 × 106 m3 mol−1 s−1 ) at 100 A m−2 : C2 H5 OH + 12OH • → 2CO2 + 9H2 O z = 12 100 s = 1.5 × 10−4 mol m−3 cHO • = √ R 96485 12 · 1.9 · 106 · 103 · 2.2 · 10−9 5. .a) Estimation of the mass transfer coefficient of HCOOH:

kd =

D at invariant δ : δ

kd,HCOOH = kd,FFC

DHCOOH DFFC = kd,HCOOH kd,FFC

1.7 · 10−9 DHCOOH = 1.2 · 10−5 = 1.9 × 10−5 m s−1 DFFC 1.1 · 10−9

b) The limiting current density for the oxidation of 0.1 M (102 mol m−3 ) HCOOH: jlim = z · F · kd,HCOOH · cHCOOH where z = 2 (HCOOH + 2OH • → CO2 + 2H2 O) jlim = z · F · kd · c = 2 · 96485 · 1.9 · 10−5 · 102 = 367 A m−2 = 0.37 kA m−2

260

ELECTROCHEMICAL OXIDATION OF ORGANIC COMPOUNDS

REFERENCES 1. S. Parsons, ed., Advanced Oxidation Processes for Water and Wastewater Treatment , IWA Publishing, London, 2004. 2. C. Comninellis, Electrochim. Acta 1994, 39 , 1857–1862. 3. G. F´oti, C. Comninellis, in Modern Aspects of Electrochemistry (R.E. White, B.E. Conway, C.G. Vayenas, M.E. Gamboa-Adelco, eds.), Modern Aspects of Electrochemistry, No. 37, Kluwer Academic / Plenum Publishers, New York, 2004, pp. 87–130. 4. A. Kapałka, G. F´oti, C. Comninellis, J. Appl. Electrochem. 2008, 38 , 7–16. 5. S. Fierro, T. Nagel, H. Baltruschat, C. Comninellis, Electrochem. Commun. 2007, 9 , 1969–1974, Electrochem. Solid State Lett ., 2008, 11 , E20–E23. 6. S. Fierro, E. Passas-Lagos, E. Chatzisymeon, D. Mantzavinos, C. Comninellis, Electrochem. Commun. 2009, 11 , 1358–1361. 7. B. Marselli, J. Garcia-Gomez, P.A. Michaud, M.A. Rodrigo, C. Comninellis, J. Electrochem. Soc. 2003, 150 , D79–D83. 8. G. Liu, J. Zhao, H. Hidaka, J. Photochem. Photobiol. A: Chem. 2000, 133 , 83–88. 9. G.R. Buettner, Free Rad. Biol. Med . 1987, 3 , 259–303. 10. M. Hermes-Lima, N.C. Santos, J. Yan, M. Andrews, H.M. Schulman, P. Ponka, Biochim. Biophys. Acta 1999, 1426 , 475–482. 11. NDRL Radiation Chemistry Data Center, http://allen.rad.nd.edu. 12. P.A. Michaud, Comportement anodique du diamant synth´etique dope au bore, PhD thesis, EPFL, No 2595, 2000. 13. J.P. Hoare, in Standard Potentials in Aqueous Solution (A.J. Bard, R. Parsons, J. Jordan, eds.), Marcel Dekker Inc., New York, 1985. 14. A. Fujishima, Y. Einaga, T.N. Rao, D.A. Tryk, eds., Diamond Electrochemistry, BKC Inc., Tokyo & Elsevier B.V., Amsterdam, 2005. 15. A. Kapałka, G. F´oti, C. Comninellis, Electrochim. Acta. 2009, 54 , 2018–2023. 16. A. Kapałka, B. Lanova, H. Baltruschat, G. F´oti, C. Comninellis, Electrochem. Commun. 2008, 10 , 1215–1218. 17. R. Scholes, J. Weiss, Radiat. Res. 1959, 1 , 177–189. 18. G.R.A. Johnson, J. Weiss, Chem. & Ind . 1955, 358–359. 19. W.M. Garrison, H.R. Haymond, W. Bennett, S. Cole, Radiat. Res. 1959, 10 , 273–282.

11 Modeling of Electrochemical Process for Water Treatment Using Diamond Films Onofrio Scialdone and Alessandro Galia

11.1

INTRODUCTION

In electrochemistry, as in other chemical methods, situations in which a single reactive path is obtained are seldom encountered. In most experimental cases the target reactive path is accompanied by side reactions leading to unwanted side products and/or to a higher energetic consume. Therefore, the optimization of the process becomes obviously a relevant issue. Quite often the high number of operative parameters that may be adjusted makes an empirical investigation exceedingly onerous in order to individuate the conditions that allow the optimization of an electrochemical process. In this perspective, theoretical models can offer useful strategies for both the individuation of the parameters that affect the selectivity of the process, and the design of new materials/apparatuses that can favor the selective occurrence of the desired route [1]. Furthermore, the experimental validation of mathematical models can furnish precious indications for scale-up stages and confirm the assumptions on which the model is based, thus allowing a proper description of the process. In the field of the electrochemical abatement of organic pollutants in wastewater, unwanted reactions should be severely minimized in order to avoid the formation of secondary pollutants and/or an increase of energetic costs. Hence, the Synthetic Diamond Films: Preparation, Electrochemistry, Characterization, and Applications, First Edition. Edited by Enric Brillas and Carlos Alberto Mart´ınez-Huitle.  2011 John Wiley & Sons, Inc. Published 2011 by John Wiley & Sons, Inc.

261

262

MODELING OF ELECTROCHEMICAL PROCESS FOR WATER TREATMENT USING DIAMOND FILMS

modeling of these processes has attracted the attention of numerous researchers in the last several years [2–10]. Both oxidation and reduction routes were used for the treatment of organic compounds in water [11,12]. Oxidation processes allow the conversion of organic pollutants in carbon dioxide (electrochemical “incineration” or “combustion”) (see Equation 11.1), to nontoxic compounds or to biocompatible organics that can be treated in a conventional biological process. Combined processes are more difficult to optimize. Therefore, the development of an oxidation process that allows treating the wastewater without the necessity of a posttreatment stage is more often pursued by researchers. Reduction processes are usually used for the conversion of some organic pollutants to nontoxic compounds. For example, halogenated aliphatic compounds can be electrochemically reduced to corresponding dehalogenated aliphatic hydrocarbons (see Equation 11.2 for the conversion of chloroalkanes to corresponding alkanes) [12], whereas nitrate ions can be converted in nitrogenous gas [13]. Cm Hn + 2mH2 O → mCO2 + (n + 4m)H+ + (n + 4m)e−

(11.1)

Cm H2m+2−p Clp + 2pe− + p H+ → Cm H2m+2 + p Cl−

(11.2)

Focusing on the oxidation route, it is relevant to observe that many organic pollutants can be anodically oxidized at a lower potential with respect to that involved for oxygen evolution, but in these conditions a decrease of the anode activity is often observed as a consequence of poisons formation [2]. These poisoning species can usually be oxidized only at high anodic potentials in the region of oxygen evolution. Therefore, the anodic abatement of organic pollutants is generally carried out at high potentials with simultaneous oxygen evolution. At these potentials the oxidation of organics can take place both by a direct anodic oxidation or by means of hydroxyl radicals generated by water oxidation. Depending on the nature of the electrode material, hydroxyl radicals can be present as free radicals in a reaction layer adjacent to the electrodic surface, with a thickness drastically lower than that of the diffusion layer due to the fast chemical evolution of these very reactive species, or adsorbed to the anodic surface. Both a weak physical and a chemical adsorption are furthermore possible depending on the electrodic material. A simplified mechanism for the electrochemical selective oxidation or combustion of organics with simultaneous oxygen evolution was proposed in 1994 by Comninellis [2]. According to this mechanism, selective oxidation to stable compounds occurs with oxide anodes (MOx ) forming a so-called higher oxide MOx +1 (chemisorbed hydroxyl radicals) and combustion occurs with phisisorbed hydroxyl radicals. On the basis of this assumption, Comninellis proposed two theoretical expressions for the instantaneous current efficiency (ICE) for the organics oxidation achieved by means of physical adsorbed hydroxyl radicals at “nonactive anodes” (electrodes which do not participate in the oxidation) or chemical adsorbed oxygen at “active anodes” (electrodes that participate in the oxidation) [2]. First, the author focused on processes that occurred in the absence of mass transfer limitations. Later, Comninellis and, co-authors (Simond, Schaller, and Comnindellis [3] and Simond and Comninellis [4]) developed a theoretical model for active electrodes that included the effect of the substrate concentration ([RH ]b ) polarization. More recently, an extension of the above mentioned theoretical models was proposed, which takes in account the occurrence of both direct anodic oxidation and oxidation mediated by physical or chemical adsorbed hydroxyl radicals [5]. Mass transfer control,

11.2 THEORETICAL MODELS

263

oxidation control, and mixed kinetic regimes were considered. Theoretical predictions were in good agreement with the experimental results obtained for oxalic and formic acids oxidation at Ti/IrO2 –Ta2 O5 anode [5,6]. In general, in the literature it has been observed more times that the weaker the interaction between hydroxyl radical with the electrodic surface, the lower is the electrochemical activity toward oxygen evolution (high oxygen overvoltage anodes) and the higher is the chemical reactivity toward organics oxidation [14]. In this frame, the high activity of Boron-doped diamond (BDD) toward the electrochemical incineration of a wide class of organics can be rationalized on the bases of the weak interactions between the diamond surface and hydroxyl radicals. These interactions are described to be so weak that the HO • can be considered as quasi-free, thus leading to high overpotential for oxygen evolution and to high reactivity toward organics oxidation [14]. The oxidation of organic pollutants can also take place in the homogeneous phase by electro-generated agents such as active chlorine, ozone, S2 O8 2− and CeIV [11]. Active chlorine or peroxidisulfuric acid, for example, can be generated by the oxidation of chlorides or sulfate ions, respectively, at the anodic surface. Particular interest has been focused on chlorine mediation, due to the ubiquitous character of Cl− species in wastewaters and its effective action. The homogeneous oxidation performed by means of active chlorine could coexist with the direct oxidation of the organics at the electrode surface, the oxidation mediated by electro-generated hydroxyl or oxychloro radicals or with both these paths. As a consequence, in this case it is not easy to evaluate a priori the effect of operative parameters on the performances of the process. In this context it highlights again the potential paramount importance of theoretical models to select the key operative parameters and design suitable materials/apparatus. So, many authors have focused their attention on the modeling of electrochemical oxidation of organic pollutants in water. Particular attention has been devoted to processes performed at BDD, due to the extreme efficacy of this electrode toward the oxidation of numerous pollutants. The modeling of both direct and indirect oxidation processes was attempted by various authors. A brief review of these studies is reported here.

11.2 11.2.1

THEORETICAL MODELS General Considerations

The electrochemical oxidation of water at BDD has been extensively investigated by various authors [11,15]. As prepared boron-doped diamond shows hydrogenated terminations that are oxidized before the oxygen evolution to oxygen-containing functionalities [16,17]. It has been suggested that hydroxyl groups are the result of strong electrochemical oxidation. Oxygen evolution at BDD starts at about 2.3 V versus SCE by oxidation of water and formation of hydroxyl radicals (see Equation 11.3) [15]. The experimental evidence for the formation of hydroxyl radicals on BDD electrodes was reported by Marselli et al. [18] by spin trapping, using electron spin resonance. Recently, Enache and co-authors have shown by DP voltammetry an oxidation peak at 2.1 V versus Ag/AgCl associated with the water discharge process and the electrochemical generation of hydroxyl radicals [19]. Many authors have supposed, on the bases of the poor adsorption ability of the BDD surface, that hydroxyl radicals generated by water oxidation at BDD are quasi-free or very weakly adsorbed on the electrode surface [14]. The

264

MODELING OF ELECTROCHEMICAL PROCESS FOR WATER TREATMENT USING DIAMOND FILMS

oxygen evolution can involve different mechanisms. In particular, two probable mechanisms involve the oxidation of the hydroxyl radical (Equations 11.4 and 11.6) or the coupling of two HO • (Equations 11.5 and 11.6) [20–22]. Note that coupling of HO • can proceed toward the formation of hydrogen peroxide (Equation 11.7), which, at its turn can be oxidized to oxygen (Equation 11.8) [21,22]. Reactions 11.4 and 11.5 can also coexist, but it is not easy, up to present knowledge, to assert what route should predominate. In this context one can consider that recombination of hydroxyl radicals is likely to be more favored on surfaces that present, in contrast from what is expected for BDD, a high active surface sites concentration [22]. On the other hand, the occurring of the recombination reaction at BDD is supported by the fact that Michaud et al. observed low H2 O2 concentrations during HClO4 electrolyses at BDD [21]. H2 O → HO • + H+ + e− HO

•

•

+

→O +H +e •

−

2O

(11.4)

•

(11.5)

→ O2

(11.6)

2HO → O + H2 O •

(11.3)

•

2HO → H2 O2

(11.7)

H2 O2 → O2 + 2H+ + 2e−

(11.8)

In the presence of organics, hydroxyl radicals are involved in their oxidation as proposed by Feng and Johnson [23], that of course takes place in competition with the oxygen evolution. BDD( • OH)n + RH → BDD + CO2 + H2 O

(11.9)

Furthermore, organics can be directly oxidized at the anodic surface or they can be oxidized by means of other electro-generated reagents such as O3 , H2 O2 , H2 S2 O8 , and so on [11]. Direct anodic oxidation and oxidation by means of hydroxyl radicals, whose main stages occur at the electrode surface, were often classified in literature as “direct processes”, whereas oxidation by means of other electro-generated oxidants, that usually take place in homogeneous phase, are called “indirect processes” [11]. Please consider that, in the case of free hydroxyl radicals, this classification is not so stringent. On the other hand, for the sake of simplicity, all the oxidation routes involving reactions with hydroxyl radicals were often grouped among the “direct processes” [11]. Concerning the “direct processes”, it is useful to observe that anodic oxidation could be potentially performed at BDD for some organics, such as phenol derivatives, at lower potentials with respect to that necessary for oxygen evolution, but the formation of a passive layer at these potentials usually prevents this possibility [7,8,24,25]. Hence, abatement of organics is usually performed in the range of potential of oxygen evolution. In these conditions, no formation of the passivation layer occurs, oxidation of organics can coexist with oxygen evolution [8–10], and both anodic oxidation and oxidation by means of hydroxyl radicals can potentially concur to the abatement of the organics.

11.2 THEORETICAL MODELS

265

Please note that in the following “direct” and “indirect” processes will be studied separately because a very different effect of operative parameters on the performances of the wastewater treatment can occur with these different groups of reactions. Amperostatic electrolyses will be considered, since the galvanostatic mode is general preferred from an applicative point of view.

11.2.2 Oxidation of Organic Pollutants in Water at BDD by Means of Direct Anodic Oxidation or Reaction with Electro-Generated Hydroxyl Radicals (‘‘Direct Processes’’) When the abatement of organic pollutants takes place by means of direct anodic oxidation or reaction with electro-generated hydroxyl radicals (“direct processes”), the oxidation processes arise on the anode surface or in a thin reaction layer adjacent to the electrode surface with a thickness dramatically lower with respect to that of the diffusion layer. Thus, hydroxyl radicals generated at BDD are expected to exist as weakly physical adsorbed species on the anodic surface or as free molecules that can diffuse through a very thin portion of the diffusion layer due to their very high reactivity. It follows that the oxidation process can be considered in these conditions as a surface or a pseudo-surface process [26] that can take place under oxidation reaction control, mass transfer control or mixed kinetic regimes depending on the rate of mass transfer of the pollutant toward the anodic surface in comparison with the oxidation rate. Please note that the process could also be under the kinetical control of adsorption or desorption stages. On the other hand, these stages are usually neglected for the poor adsorption ability of the BDD surface. When the rate of the mass transfer of the pollutant is dramatically lower than that of its anodic oxidation, the concentration of the pollutant at the anodic surface/reaction layer C0 is close to zero and the oxidation process is under mass transfer control. This case arises, for an oxidation process that proceeds up to the total oxidation of the organic pollutant, when the limiting current density ilim = nFkm [RH]b ≪ iapp ICEOC (where n is the number of electrons exchanged for the anodic oxidation of RH to carbon dioxide; F is the Faraday constant (96487 C mol−1 ); [RH]b and km are the bulk concentration and the mass transfer coefficient of the organic RH, respectively; iapp is the applied current density; and ICEOC is the instantaneous current efficiency for the oxidation of RH under oxidation reaction control under adopted operative conditions). Conversely, when ilim ≫ iapp ICEOC , mass transfer is significantly faster with respect to oxidation rate, C0 is very close to the concentration in the bulk Cb and the process is under reaction oxidation control. When more pollutants are present in the bulk during the electrolysis, one can focus on the chemical oxygen demand (COD) of the solution. In particular, the limiting current density, for a process that proceeds up to the total oxidation of organics, is given by ilim = 4Fkm COD and a mass transfer control will arise if ilim ≪ iapp ICEOC (where ICEOC is the average current efficiency for a process under oxidation reaction control and COD is the chemical oxygen demand computed on a molar base). The effect of operative parameters on the performances of the process was recently investigated [5,6,30,31], by considering, for the sake of simplicity, the complete oxidation of an organic pollutant that proceeds in competition with the oxygen evolution with no formation of intermediates with significant concentrations in the bulk. In this case, the instantaneous current efficiency for a process under galvanostatic mode can be rapidly estimated for both mass transfer and oxidation reaction control regimes.

266 •

MODELING OF ELECTROCHEMICAL PROCESS FOR WATER TREATMENT USING DIAMOND FILMS

Under mass transfer control, the current efficiency ICEMT is given by the ratio between the limiting current ilim and the applied current density iapp independently by the oxidation mechanism [5,7,8,30,31]. For [RH]b ≪ C ∗ ICEOC , ICE = ICE MT ≈ [RH]b /C ∗ = nFkm [RH]b /iapp (11.10) where C ∗ = iapp /nFkm . It follows that, under mass transfer control, the ICE should depend on the fluidodynamics of the system (trough its effect on km ), on the applied current density and on [RH]b (with a linear dependence on this parameter). It is important to observe that, under these conditions, the performances of the process are readily predictable. Thus, the mass transfer coefficient is readily given by the ratio km = D/δ where the diffusion coefficient D can be often found in literature or can be estimated by electroanalytical experiments or using the Wilke-Chang expression, whereas the thickness of the diffusion layer δ is easily estimated by typical limiting current essays using, as an example, the couple hexacyanoferrate (II)/hexacyanoferrate (III). Hence, it is also possible to predict the organics concentration as a function of the charge or of the time. Indeed, the instantaneous current efficiency is given, by definition, by: ICE = −nFV d[RH]b /dQ

(11.11)

and eliminating the term ICE by Equations 11.10 and 11.11, one easily obtains the relationships reported in Equations 11.12 and 11.13 (where V is the volume, A is the anodic surface, and t and Q are the time and the charge passed, respectively). [RH]b,Q = [RH]b,Q = 0 exp[−Q/(nFVC ∗ )] [RH]b,t = [RH]b,t = 0 exp[−(Akm /V )t] •

(11.12) (11.13)

For a process under the kinetic control of the oxidation reaction, the instantaneous current efficiency ICEOC is determined by the competition between the oxidation of the organic and the evolution of the oxygen [5,30,31,40]: For [RH]b ≫ C∗ ICEOC

iRH iRH = iapp iRH + iO2 1 1 = = iO2 [RH]∗ 1+ 1+ iRH [RH]b

ICEOC =

(11.14)

where iRH and iO2 are the current densities involved in the oxidation of the organic and in the oxygen evolution process, respectively. The term [RH]∗ is the value of [RH]b , which gives a current density for the RH oxidation iRH equal to the current density involved for the oxygen evolution reaction iO2 (e.g. the value of [RH]b that gives ICE = 50%) [5]. Please consider that the physical meaning of [RH]* depends on the oxidation route. Thus, if the oxidation of the organic takes place by direct anodic oxidation [RH]* is given by [RH]∗dir = 2r(E )/nk (E ), where k (E ) is the heterogeneous rate constant for the oxidation of RH and r(E ) is the rate of

11.2 THEORETICAL MODELS

267

the solvent oxidation. Otherwise, if the oxidation takes place by means of hydroxyl radicals, [RH]∗HO is given (as an example, considering for the sake of simplicity physically adsorbed hydroxyl radicals) by: [RH]∗OH = 2kor /nkO2

(11.15)

where kO2 (s−1 ) and kor (M−1 s−1 ) are the rate constants of Reactions 11.16 and 11.17, respectively, and n is the number of adsorbed hydroxyl radicals necessary to convert the substrate to carbon dioxide [5,7,30]. BDD( • OH) → BDD + 0.5 O2 + H+ + e−

(11.16)

BDD( • OH) + RH → BDD + H2 O + R •

(11.17)

It is interestingly to observe that [RH]∗dir assumes a constant value with the potential (and applied current density), given by the following relationship, if the transfer coefficients α of the two competitive oxidation routes assume similar values: [RH]∗dir = (2/n)(k ′ /k

◦

RH )exp[(1

− α)F (E

◦ RH

−E

◦

W )/(RT)]

(11.18)

where k′ is given by the product of the standard rate constant for the oxidation of the solvent and the solvent concentration and E ◦ W and E ◦ RH are the standard potentials for the oxidation of the water and RH, respectively [5,7,30]. Hence, in the case of a direct anodic oxidation process, expression reported in Equation 11.14 can be readily used to predict the effect of some operative parameters such as current density and organic concentration on the performances of the process. A more complex scenario is expected for an indirect process mediated by hydroxyl radicals. In this case, in fact, [RH]∗ is expected to depend on the working potential (e.g., on the applied current density) so that a less easy prediction of the effect of operative parameters on ICE is achievable [5,30]. Anyway, this aspect is of less practical relevance if the kinetic constant of mediated reaction is so high that in any case [RH]∗ is close to zero and ICE under oxidation control is close to 1. In general, it is interesting to observe that according to Equations 11.10 and 11.14, the instantaneous current efficiency of the process is expected to depend on the concentration of the organic pollutant in the bulk [RH]b for both mass transfer and oxidation reaction control. On the other hand, in the case of a mass transfer control, ICE should depend linearly on [RH]b , whereas in the case of an oxidation reaction control, a linear relationship is expected between 1/ICE and 1/[RH]b . Furthermore, if the kinetic constant of the reaction between the organic pollutant and hydroxyl radical is very high, as usually expected for diamond anodes, for a process under oxidation reaction control, the dependence of the ICE by [RH]b can be practically observed only for very low values of both [RH]b and the applied current density (e.g., by working with electrodes characterized by a very large surface area). Please note also that, under oxidation reaction control, the process is expected to be dramatically affected by the nature of both the organic pollutant and of the electrode material through their effect on [RH]* but not by the flow dynamic regime. However, in the case of a mass transfer control, the ICE depends strongly on the flow dynamic regime (trough its effect on km ), slightly on the nature of the organic pollutant (trough the value of the diffusion coefficient) but not significantly on the nature of the anode. It follows

268

MODELING OF ELECTROCHEMICAL PROCESS FOR WATER TREATMENT USING DIAMOND FILMS

that the utilization of boron-doped diamond anodes is expected to affect positively the abatement of organic pollutants in water if the process is under oxidation reaction control but not if a mass transfer control arises. On the other hand, please consider that a mass transfer control is expected when [RH]b ≪ C∗ ICEOC . Therefore, if the same pollutant is treated at diamond anode with a high value of ICEOC (e.g., a low value of [RH]*) or at an electrode characterized by a very low value of ICEOC , for the same operative conditions BDD can experience a mass transfer control and the other electrode an oxidation reaction control with a slower rate with respect to that of the mass transfer at BDD. This case, as an example, was observed for the oxidation of oxalic acid in basic conditions at BDD and Ti/IrO2 -Ta2 O5 [6,42]. Thus, the oxidation reaction was so slow at the iridium-based anode that a mass transfer control never took place at this electrode under adopted operative conditions [42]. Let us now discuss the general case of a process whose rate determining step changes during the electrolysis from oxidation reaction control in the first stages to mass transfer control in the last part with a mixed regime between them. In this case, a general expression for the instantaneous current efficiency has to be adopted. In particular, under the previously mentioned conditions, it is easy to show that the following expression applies [5,30,31,40]: ICE =

1 2[RH]∗ 1+ [RH]′ + ([RH]′ 2 + 4[RH]∗ [RH]b )0.5

(11.19)

where [RH]′ = [RH]b − [RH]*−C * and C ∗ = i /nFkm . Theoretical predictions based on Equations 11.10, 11.14, and 11.19 were used, up to now, to predict the effect of operative parameters on the abatement of two quite simple molecules such as oxalic and formic acid. A very good agreement between experimental data and theoretical predictions was observed changing severely the current density, the flow rate, and the initial concentration of the acid (e.g., see Figures 11.1 and 11.2 for the oxidation of oxalic acid). This result appeared rather interesting also considering that the kinetic constant of the reaction between hydroxyl radicals and the organic presents very different values for these two acids. Please note that numerous pollutants can often be present in the water system from the beginning of the electrolysis or as a result of the formation of some intermediates with appreciable concentrations in the bulk. In this case, the modeling of the process has to take in account the contemporaneous presence of various organic pollutants in the solution. This general problem was studied by various groups, including those of Comninellis, Polcaro, Rodrigo, and Scialdone with different approaches. Some of these studies are briefly summarized in the following. 11.2.2.1 The Model of Comninellis and Coauthors Comninellis and coauthors observed that the oxidation of many organics, such as phenol, proceeds at BDD toward the complete incineration with very high current efficiency if no mass transfer limitations arise [7,8,24,25]. On the bases of these results, the authors developed a very simple kinetic model for a batch recirculation system with galvanostatic alimentation based on the assumption that when the oxidation of organics is performed at BDD at high anodic potentials, close to oxygen evolution, the electrochemical incineration of the organic

11.2 THEORETICAL MODELS

269

(a) 0.12

[OA]b(M)

0.10 0.08 0.06 0.04 0.02 0.00 0

2000

4000 Q(C)

6000

8000

0

2000

4000 Q(C)

6000

8000

0

2000

4000 Q(C)

6000

8000

(b) 0.12 0.10

[OA]b(M)

0.08 0.06 0.04 0.02 0.00

(c) 0.12 0.10

[OA]b(M)

0.08 0.06 0.04 0.02 0.00

Figure 11.1 Profile of oxalic acid concentration [OA]b versus charge passed for electrolysis performed at BDD with (a) 17, (b) 39, and (c) 56 mA/cm2 , at different flow rates. Flow rate: 1.7 (), 3.4, (-) 4.9 l/min (•). Amperostatic electrolyses. Theoretical curves (—) obtained by Equation 11.19 with [RH]* = 0.013 M. System solvent supporting electrolyte (SSE): Water, Na2 SO4 , H2 SO4 . Initial oxalic acid concentration: 100 mM. T = 25◦ C. (Reprinted with permission from Ref. 30.)

270

MODELING OF ELECTROCHEMICAL PROCESS FOR WATER TREATMENT USING DIAMOND FILMS

(a) 0.20 0.04 0.15 [OA]b(M)

0.02 0.00

0.10

0

5000

10000

15000

0.05

0.00 0

5000 10000 Charge passed (C)

15000

0

0.05

0.15

(b) 1.2 1

ICE

0.8 0.6 0.4 0.2 0

0.1 [OA]b(M)

(c) 5

1/(ICE)

4

3

2

1

0 0

40 ([OA]b)–1(M–1)

80

Figure 11.2 Oxalic acid concentration [OA]b versus (a) charge passed, (b) ICE versus [OA]b , and (c), 1/ICE versus 1/[OA]b . Initial substrate concentration: 0.01 (-), 0.05 (•), 0.1 () and 0.2 M (◦). Flow rate: 4.9 l/min. Current density: 17 mA/cm2 . Amperostatic electrolyses at BDD anode. System solvent supporting electrolyte (SSE): Water, Na2 SO4 , H2 SO4 . T = 25◦ C. Theoretical curves (—) in Figure 2a obtained by Equation 11.19 with [RH]* = 0.013 M. Theoretical curves in Figures 2b and 2c (–) obtained by Equations 11.14 and 11.10 with [RH]* = 0.013 M. (Reprinted with permission from Ref. 30.)

11.2 THEORETICAL MODELS

271

compound is a fast reaction and it is controlled by mass transport toward the anode. Particularly, authors assumed that: The reservoir volume (V ) is much greater than that of the electrochemical reactor. The electrochemical reactor and the reservoir are perfectly mixed. • The oxidation in the bulk by electro-generated oxidants is not considered. • If iapp < ilim (e.g., if the electrolysis is under current limit contro)l, the current efficiency for the abatement of the organic is 100%. As a consequence, COD decreases linearly with time, as shown by the relationship reported in the following equation (where COD◦ is the initial COD, A is the anodic surface, V is the solution volume, and t is the time). •

•

  iapp A α ′ km A ◦ t = COD 1 − t COD = COD − nFV V ◦

(11.20)

◦ where α ′ = iapp /ilim . • If iapp > ilim (e.g., if the electrolysis is under mass transfer control), secondary reactions (such as oxygen evolution) commence, resulting in a decrease of COD. The current efficiency can be simply estimated as

ICE = ICEMT ≈ ilim /iapp

(11.21)

and, as a consequence, the COD removal follows an exponential trend described by the following equation (where CODcr and tcr are the initial values of COD and time when the process is under mass transfer control from the beginning of the electrolysis or the values of COD and time achieved when iapp = ilim ):     1 − α′ Akm Akm ◦ ′ (t − tcr ) = α COD exp − t+ COD = CODcr exp − V V α′ (11.22) It is important to underline the fact that this very simple model does not present any adjustable parameter. Thus, the mass transfer coefficient is readily given by the ratio km = D/δ, where the diffusion coefficient D can be often found in literature or can be estimated by electroanalytical experiments or using the Wilke-Chang expression, whereas the thickness of the diffusion layer δ is easily estimated by typical limiting current essays using, as an example, the couple hexacyanoferrate (II)/hexacyanoferrate (III). Of course, this model is expected to fit with accuracy experimental data when the current efficiency in the absence of mass transfer limitations is close to 1. Furthermore, the model does not describe mixed kinetic regimes or the evolution of the concentrations of different species present in the system. On the other hand, this simple model presents a very good agreement with experimental data for several organic compounds and various operative conditions for the evolution of COD with the time or the charge passed [7,8,14,15,24,25]. (see, as an example, Figures 11.3 and 11.4 for the case of the oxidation of 2-naphtol at acidic pH.) Therefore, it represents a very interesting tool for the prediction of the effect of operative parameters, such as current density, flow rate, and organic concentration, on the “direct” abatement of COD at a BDD anode in a batch recirculation system.

272

MODELING OF ELECTROCHEMICAL PROCESS FOR WATER TREATMENT USING DIAMOND FILMS

Figure 11.3 Influence of the initial 2-naphthol concentration () 9 mM; ( × ) 5 mM; (•) 2 mM on the trends of COD and ICE (inset) during electrolysis, using a BDD anode. Supporting electrolyte: 1 M H2 SO4 ; T = 30◦ C; j = 30 mA cm−2 . The solid lines represent the model prediction. (Reprinted with permission from Ref. 8.)

60

1.2 1 0.8 ICE

COD (mol O2m–3)

50 40

0.6 0.4 0.2

30

0 0

2

4

6

8

10

14

16

Q (Ah dm–3)

20 10 0

0

2

4

6

8 10 Q (Ah dm–3)

12

Figure 11.4 Influence of applied current density (•) j = 15 mA cm−2 ; (×) j = 30 mA cm−2 ; () j = 60 mA cm−2 on the trends of COD and ICE (inset) during electrolysis in 1 M H2 SO4 + 5 mM 2naphthol, using a BDD anode. T = 30◦ C. The solid lines represent the model prediction. (Reprinted with permission from Ref. 8.)

11.2.2.2 The Theoretical Works of Polcaro and Coauthors Polcaro and coauthors reported a more complex model with the aim of evaluating the trends of the concentrations of the starting pollutant and intermediate products in the stagnant layer and in the bulk [26]. Thus, it has been shown that the incineration of some organics at BDD proceeds trough the formation of intermediates. As an example, the oxidation of

11.2 THEORETICAL MODELS

273

phenol can be represented by the following schematic reaction path: Phenol → Cyclic intermediates → Aliphatic acids → CO2 Furthermore, this model applies also to the case of an oxidation process that gives rise to a current efficiency lower than 100% in the absence of mass transfer limitations. The model was based on the following assumptions: Oxidation of pollutants and intermediate compounds takes place by homogeneous chemical reactions with hydroxyl radicals (with first-order kinetic reaction with respect to both organics and hydroxyl radicals concentrations) in the diffusion layer in competition with the chemical deactivation of hydroxyl radicals, which is described by a first-order kinetic reaction not depending on the working potential. • A diffusion-reaction model is used to model the diffusion layer, whereas the bulk of the solution is represented by a stirred-tank reactor. •

The model required a numerical solution of pertaining equations (mass balance equations in the bulk and in the stagnant layer for the starting pollutant and intermediates and in the diffusion layer for hydroxyl radicals). As a result, authors were able to predict the trend with time of the concentration of different compounds present in the bulk of the solution during the electrolyses, as well as the evolution of the space profile of the species. As an example, the authors compared the theoretical predictions of the model with experimental data for the oxidation of Phenol at BDD. The model was able to predict with a good accuracy the trend with time of the concentrations of phenol and main intermediates (aromatic and aliphatic acids) by changing significantly the current density and the flow dynamic regime, and using one adjustable parameter (the average kinetic constant between hydroxyl radicals and aromatic compounds) [26]. A good agreement between theoretical predictions and experimental data was reported by the authors also in the cases of cyanuric acid and atrazine by using, as an adjustable parameter, the kinetic constant between hydroxyl radicals and the adopted pollutant [26]. As mentioned, the model involved also the calculation of the concentration profiles in the diffusion layer of hydroxyl radicals, pollutants, and intermediates with the time. Quite interestingly, according to this model, hydroxyl radicals diffuse in the stagnant layer for few tens of nanometers. As an example, a thickness of the reaction layer of about 50 nm was computed for a current density of 25 mA/cm2. 11.2.2.3 The Approach Proposed by Rodrigo and Coauthors In order to reduce the mathematic complexity of theoretical models aimed to determine the concentrations of compounds involved in the oxidation process, Rodrigo and coauthors considered some simplifying assumptions. Their model divides the electrochemical reactor into three zones: two are close to the electrodes (electrochemical zones) and a third zone corresponds to the bulk of the solution (chemical zone) that is considered as consecutive stirred-tank reactors [27–29]. Hence, in each zone, the concentration is assumed to depend on the time passed but not on the position. The concentration of each compound in the chemical zone is taken as the value measured experimentally. Mass transport processes between the electrochemical and chemical zones are quantified by assuming that the local exchange rate is proportional to the difference of concentration between the two zones. The authors observe that a number of processes can occur at the electrode

274

MODELING OF ELECTROCHEMICAL PROCESS FOR WATER TREATMENT USING DIAMOND FILMS

surface, and therefore the total current applied is shared among all these processes. The rate of each process is given by ri = (iapp A/F)αielectrode

(11.23)

where αielectrode , which gives the ratio of current density for a particular electrochemical process with respect to the applied current density, is a measure of the relative oxidizability of a compound and αielectrode = 1. In a first version of the model, αielectrode was a function of the oxidizability factors of the compounds present in the system that were computed as adjustable parameters [27]. Later, the authors assumed that the fraction of iapp used in each process depends on the cell potential (Vwork ) and the oxidation (or reduction) potential (Vi ) of each process by the relationship reported below, so that no adjustable parameters were used in the model [28]. αielectrode = (Vwork − Vi )/



(Vwork − Vi )

(11.24)

The model accounted also for the presence of homogeneous oxidation reactions between electro-generated oxidants and organics assuming second-order rate expressions depending on the concentrations of both the oxidant and the organic compound. In the case of phenol and carboxylic acids, no presence in the bulk of oxidants was evaluated with a I2 /I− test [28]. It was therefore concluded that no mediated oxidation processes occurred in the bulk zone. Hence, mediated process, involving oxidation by means of hydroxyl radicals, might occur only in the electrochemical zone and were considered as direct reactions. The model showed a good agreement with a very large set of experimental data obtained for the incineration of various organic compounds, including phenol and different carboxylic acids at BDD in a water solution of Na2 SO4 [28]. In the case of phenol, for example, the model considered the direct oxidation of phenol at the anode to carbon dioxide, through the formation of maleic and oxalic acids as intermediates, and the hydrogen evolution at the cathode (see Figure 11.5). As shown in Figure 11.6, theoretical predictions were in a good agreement with experimental data obtained under a very large range of operative conditions.

11.2.3 Oxidation of Organic Pollutants in Water by Means of Electro-Generated Oxidants (‘‘Indirect Processes’’) Such as Active Chlorine As previously mentioned, electrochemical oxidation of organics can occur also by the action of electro-generated oxidants such as active chlorine, ozone, peroxidisulfuric acid, CeIV , and more. Hence, various researchers included in their model the homogeneous oxidation of organics by means of electro-generated oxidants [10,28,29,40,44]. It is interesting to observe that in indirect processes the oxidation of the organics take place mainly by a homogeneous reaction and a very different effect of operative parameters on the performances of the process can arise with respect to direct ones. Thus, the current efficiency of the process is expected to be determined mainly by the current efficiency due to the electro-chemical formation of oxidants and by the competition between the homogeneous chemical oxidation of organic pollutants and competitive reaction paths

11.2 THEORETICAL MODELS

Anodic reaction zone (a)

anode

1

Phenol*

e–

Cathodic reaction zone (c)

cathode

Phenol

2 3

Maleic*

e–

Chemical reaction zone (b)

275

Maleic

9

Oxalic

Oxalic*

5

H2 e–

4 H 2O

6

CO2

CO2*

7

H2O e–

8 O2

(1) r1 = k·A·([Phenol]b–[Phenol]a) (2)

anode r2 = I a2 F

(3) r3 = k·A·([Maleic]b–[Maleic]a)

(4) r4 = k·A·([Oxalic]b–[Oxalic]a)

(7) r7 = k·A·([CO2]b–[CO2]a)

(5)

anode r5 = I a5 F

I anode (8) r8 = a8 F

(6)

anode r6 = I a6 F

I cathode (9) r9 = a9 F

Figure 11.5 Sketch representing the processes considered in the modeling of the electro-oxidation of phenol-polluted wastewater using BDD anodes. (Reprinted with permission from Ref. 28.)

involving the electro-generated oxidants. In the following, this aspect will be studied in the frame of the oxidation of organics by means of active chlorine. This process was considered as a model case for numerous reasons. First, the effect of chloride ions on the performances of the process has been the object of numerous researches [11] due to the ubiquitous character of Cl− species in waste waters and to the fact that chloride ions can cause an increase in the removal efficiency of organic pollutants for the involvement of active chlorine in the oxidation process. Furthermore, it has been shown that the addition of chloride ions can give rise in some cases to the formation of halogenated intermediates more toxic than the starting compounds [32]. Second, in the presence of Cl− , a quite intriguing change of the effect of some operative parameters on the performances of the process can occur with respect to “direct processes.” When chloride ions are added to a water solution, their oxidation can lead to the formation of chlorine, hypochlorous acid, and/or hypochlorite, depending on the pH (Equations 11.25–11.27) that can oxidize the organics near to the anode or/and in the bulk of the solution (Equation 11.28) in alkaline medium [33–36]. 2Cl− → Cl2 + 2e− Cl2 + H2 O → HOCl + H+ + Cl−

(11.25) (11.26)

35

Concentration (mmol C dm–3)

MODELING OF ELECTROCHEMICAL PROCESS FOR WATER TREATMENT USING DIAMOND FILMS

50

Concentration (mmol C dm–3)

Concentration (mmol C dm–3)

276

(a)

60 50 40 30 20 10 0 0

5

10

15

20

25

30

60

(b)

50 40 30 20 10 0 0

5

Concentration (mmol C dm–3)

Charge (A h dm–3) (c)

80 70 60 50 40 30 20 10 0 0

10

20

30

40

Charge (A h dm )

25

(d)

70 60 50 40 30 20 10 0

–3

10 15 20 Charge (A h dm–3)

0

10

20 30 40 50 Charge (A h dm–3)

60

70

Figure 11.6 Results obtained [simulation (lines) versus experimental data (points)] for four experimental runs for the abatent of phenol: () phenol, () maleic acid, () oxalic acid, (-) carbon dioxide. (a) j = 30 mA cm−2, T = 25◦ C, pH = 2. (b) j = 30 mA cm−2 , T = 25◦ C, pH = 12. (c) j = 30 mA cm−2 , T = 25◦ C, pH = 9. (d) j = 60 mA cm−2 , T = 25◦ C, pH = 12. (Reprinted with permission from Ref. 28.)

HOCl ↔ H+ + OCl− Organics + OCl− → intermediates → CO2 + Cl− + H2 O

(11.27) (11.28)

These reactions take place in competition with oxygen evolution, chlorate chemical and electrochemical formation, and cathodic reduction of oxidants in the presence of undivided cells. Amperostatic electrolyses of chlorides gives rise at DSA anodes mainly to hypochlorite and chlorate, whereas at BDD, a more complex mixture—including hypochlorite, chlorine dioxide, and reactive oxygen species—was found [32,37]. The authors supposed that these species could be formed by reaction paths involving the reaction between hydroxyl radicals and active chlorine. It has been proposed [33,38] that adsorbed chloro- and oxychloro-radicals could be also involved in the oxidation mechanism. Furthermore, the possibility that some role could be played by the anodic shift of the oxygen evolution, caused by Cl− ions in the solution, has also been taken in consideration for platinum electrodes [38,39]. It follows that the oxidation of organics performed by means of active chlorine could coexist with the direct oxidation at the electrode surface, the reaction with hydroxyl or oxychloro radicals, or with both these paths. As a consequence, it is not easy for these complex systems to rationalize the effect of operative parameters, on the performances of the process. Furthermore, the coexistence of these oxidative routes makes this process completely different from the chemical oxidation with hypochlorite. This has been confirmed by many experimental works that usually report higher abatement of organic pollutants for the electrochemical oxidation with chlorides with respect to the chemical one with hypochlorite [33–38].

11.2 THEORETICAL MODELS

277

A very simplified theoretical approach was recently presented with the aim to examine how the more important operative parameters are expected to affect the performances of the process both in the absence and in the presence of chlorides [40]. In particular, it was assumed, for the sake of simplicity, that the reaction between the organics and active chlorine takes place mainly in the bulk of the solution. In the absence of chlorides or in the presence of a very high ratio between the concentration of organics and NaCl, main processes should be “direct” ones, such as a direct anodic oxidation or a mediated oxidation by hydroxyl or oxychloro radicals. Hence, the current efficiency for the abatement of the organics should be favored, as previously explained in detail, by (1) high values of the organic concentration; (2) the use of an electrodic material such as BDD, which provides an high oxygen over potential (e.g., a low value of k ′ (E ) and [RH]*); (3) higher flow rates and lower current densities (when the process is not under oxidation reaction control). Let us focus our attention on the opposite case of a sufficiently high ratio between the concentrations of chlorides and organics so that the contribution of the direct processes can be neglected. In this case, the current efficiency should depend mainly on the formation in the bulk of active chlorine ICE AC and on the competition between homogeneous oxidation of organics and chlorate formation [40]. Interestingly, ICEAC has been shown to increase for both BDD and DSA anodes in the presence of higher chloride concentrations and high current density, whereas a more complicated effect of the flow rate is reported. Higher and lower current efficiencies are reported by increasing the flow rate at BDD and DSA, respectively [37,40]. Hence, an opposite effect of current density is expected if the oxidation of organics takes place by a direct anodic process or by an oxidation mediated by active chlorine. Furthermore, BDD presents quite low current efficiency for the active chlorine formation with respect to low oxygen overpotential electrodes such as iridium and ruthenium oxides. Thus, the addition of chlorides in the case of BDD can result, depending on the adopted operative conditions, in a lower abatement of organics. Please note that when the direct process is under mass transfer control, the addition of chlorides is expected to result, if sufficiently high values of current density are imposed, in higher current efficiency for the mediated process. This is because in this case the process is no longer affected negatively by the kinetic limitations imposed by the mass transfer of the pollutant toward the anodic surface. Another important consequence is that BDD, which presents generally drastic higher abatement of organics with respect to DSA, can present, in the presence of suitable amounts of NaCl, a lower abatement with respect to these anodes. Real systems are complicated by the fact that direct and indirect processes could coexist under the same operative conditions. In this case, flow rate, current density, and anodic material should have an opposite effect on homogeneous and heterogeneous oxidation processes so that their overall effect on the performances of the process depends on the relative rates of heterogeneous and homogeneous oxidation reactions. Interestingly, the previously mentioned considerations on the effect of various operative parameters were confirmed by various studies [40,43,47], including one on the electrochemical incineration of oxalic acid at BDD and iridium-based anodes performed in the absence and in the presence of various amounts of NaCl [40]. Thus, a very different influence of the nature of the anodic material, the flow rate, and the current density on the performances of the process was observed in the absence and in the presence of chlorides so that optimization of the two processes required very different operative conditions. In the absence of chlorides, high current efficiency (CE) were obtained at BDD

278

MODELING OF ELECTROCHEMICAL PROCESS FOR WATER TREATMENT USING DIAMOND FILMS

when most of the process was under oxidation reaction kinetic control (i.e., when low current densities and high flow rates were imposed). On the other hand, in the presence of high concentrations of NaCl, higher CE was usually obtained at DSA with respect to BDD and at higher values of current densities. A simple theoretical model for the oxidation of organics in the presence of Cl− was proposed by Szpyrkowicz and Radaelli [44]. These authors described the kinetics of decolorization of simulated textile wastewaters in the presence of Cl− by means of a second-order rate constant: d(abs)/dt = −k (abs)[Cl2 ]

(11.29)

where (abs) is the measured absorbency and [Cl2 ] is the concentration of dissolved chlorine. The model was used by the authors, assuming that the Faradaic efficiency for the chlorine evolution is constant trough the electrolysis, with good results in the frame of the scale up of the oxidation of a textile wastewater at Ti/Pt-Ir anode from a batch to a continuous flow unit [44]. Recently, Anglada, Urtiaga, and Ortiz [45] studied in detail the electro-oxidation of landfill leachate at BDD on a pilot scale. Organic matter and ammonia oxidation were highly influenced by the applied current density value. Interestingly, the abatement of the organic matter was successfully modeled by using the approach proposed by Comninellis and coauthors (see Section 11.2.2) when a current density of 300 and 450 A m−2 was applied [45]. At higher current densities, the oxidation rate was higher than that predicted by the model, thus suggesting the occurrence of a mediated oxidation process by means of active chlorine. The authors also used the model proposed by Szpyrkowicz and Radelli [44] to predict the evolution of ammonium based on the hypothesis that the degradation takes place by a second-order reaction between ammonium and active chlorine. The kinetic model was able to predict the evolution of the concentration of the pollutant as a function of the current density in the range 300–600 A m−2 . At higher current densities a worse correlation between predicted and experimental results occurred. In these conditions a rapid decrease in chloride concentration occurred, thus suggesting that one of the model’s assumptions, the constancy of the Faradaic efficiency for the chlorine evolution during the electrolysis, was not complied [45]. It is relevant to observe that, as stated by Anglada et al. [45,46], many of the theoretical models proposed for the description of the electrochemical oxidation of organics in water still have to be validated for real wastewaters with complex and unknown detailed composition. In this case, in fact, the simultaneous occurrence of many competitive reactions at the anodic surface and in the homogeneous phase can create very complex scenarios that are very difficult to model in a simple and accurate way.

11.3

CONCLUSIONS

In the last years, several studies have proven that conducting diamond thin films can be used as very effective electrodes for the complete oxidation of organic pollutants in water and for water disinfection. The kinetic modeling of the electrochemical abatement of organics in water at diamond anodes has been attempted by many authors with the main objective of predicting the trend of the concentrations of pollutants and intermediates, of chemical oxygen demand, and of current efficiency by changing severely the operative

REFERENCES

279

conditions. The electrochemical oxidation of organic pollutants in water at diamond electrodes can take place by means of hydroxyl radicals generated by water oxidation and by direct anodic oxidation (“direct processes”) or in homogeneous phase by means of electro-generated reagents such as O3 , H2 O2 , H2 S2 O8 , active chlorine, and so on (“indirect processes”). “Direct processes” can be considered as surface or pseudo-surface processes and are strongly affected by the mass transfer rate of the organic pollutants toward the anodic surface. The modeling of these processes has been performed by various research groups, and a very good agreement between experimental data and theoretical predictions was generally reported. In “indirect processes”, the oxidation of the organics involves a homogeneous reaction with electro-generated oxidants, and a quite different effect of operative parameters on the performances of the process can take place with respect to the case of “direct processes”. Thus, in the case of “indirect processes”, the abatement of the organics is expected to depend mainly on the current efficiency due to the electrochemical formation of oxidants and on the competition between homogeneous chemical oxidation of organic pollutants and competitive reaction paths involving the electro-generated oxidants. The oxidation of organics in the presence of active chlorine was, in particular, briefly discussed as a model case.

11.4

ACKNOWLEDGMENTS

Universit`a di Palermo and Ministero dell’Istruzione, dell’Universit`a e della Ricerca (MIUR) are acknowledged for their financial support.

REFERENCES 1. C. Amatore, “Micro and Macropenomena,” in Organic Electrochemistry (H. Lund, M. Baizer, eds.), Marcel Dekker, Inc., New York, 1991, p. 220. 2. C. Comninellis, Electrochim. Acta 1994, 39 , 1857–1862. 3. O. Simond, V. Schaller, Ch. Comninellis, Electrochim. Acta 1997, 42 , 2009–2012. 4. O. Simond, Ch. Comninellis, Electrochim. Acta 1997, 42 , 2013–2018. 5. O. Scialdone, Electrochim. Acta 2009, 54 , 6140–6147. 6. O. Scialdone, S. Randazzo, A. Galia, G. Filardo, Electrochim. Acta 2009, 54 , 1210–1217. 7. M.A. Rodrigo, P.A. Michaud, I. Duo, M. Pamizza, G. Cerisola, Ch. Comninellis, J. Electrochem. Soc. 2001, 148 , D60–64. 8. M. Panizza, P.A. Michaud, G. Cerisola, Ch. Comninellis, J. Electroanal. Chem. 2001, 507 , 206–214. 9. A.M. Polcaro, S. Palmas, Ind. Eng. Chem. Res. 1997, 36 , 1791–1798. 10. A.M. Polcaro, M. Mascia, S. Palmas, A. Vacca, Ind. Eng. Chem. Res. 2002, 41 , 2874–2881. 11. C.A. Martinez-Huitle, S. Ferro. Chem. Soc. Rev . 2006, 35 , 1324–1340. 12. O. Scialdone, A. Galia, G. Filardo, Electrochim. Acta 2008, 53 , 7220–7225. 13. B.P. Dash, S. Chaudhari, Water Res. 2005, 39 , 4065–4072. 14. A. Kapalka, G. Foti, Ch. Comninellis, J. Appl. Electrochem. 2008, 38 , 7–16. 15. M. Panizza, G. Cerisola, Electrochim. Acta 2005, 51 , 191–199. 16. R. DeClements, G.M. Swain, T. Dallas, M.W. Holtz, R.D. Herric, J.L. Stickney, Langmuir 1996, 12 , 6578–6586.

280

MODELING OF ELECTROCHEMICAL PROCESS FOR WATER TREATMENT USING DIAMOND FILMS

17. I. Yagi, H. Notsu, T. Kondo, D.A. Tryk, A. Fujishima, J. Electroanal. Chem. 1999, 473 , 173–178. 18. B. Marselli, J. Garcia-Gomez, P.A. Michaud, M.A. Rodrigo, C. Comninellis, J. Electrochem. Soc. 2003, 150 , D79–D83. 19. T.A. Enache, A.M. Chiorces-Paquim, O. Fatibello-Filho, A.M. Oliveira-Brett, Electrochem. Comm. 2009, 11 , 1342–1345. 20. K. Kinoshita, in Electrochemical Oxygen Technology, The Electrochemical Society, New York,Wiley, 1992, chap. 2. 21. P.A. Michaud, M. Panizza, L. Quattara, T. Diaco, G. Foti, Ch. Comninellis, J. Appl. Electrochem. 2003, 33 , 151–154. 22. M.H.P. Santana, L.A.D. Faria, J.F.C. Boodts, Electrochim. Acta 2005, 50 , 2017–2027. 23. J. Feng, D.C. Johnson, J. Electrochem. Soc. 1990, 137 , 507–510. 24. J. Iniesta, P.A. Michaud, M. Panizza, G. Cerisola, A. Aldaz, Ch. Comninellis, Electrochim. Acta 2001, 46 , 3573–3578. 25. F. Montilla, P.A. Michaud, E. Morallon, J.L. Vazquez, Ch. Comninellis, Electrochim. Acta 2002, 47 , 3509–3513. 26. A.M. Polcaro, A. Vacca, S. Palmas, M. Mascia, J. Appl. Electrochem. 2003, 33 , 885–893. 27. P. Canizares, M. Diaz, J.A. Dominguez, J. Garcia-Gomez, M.A. Rodrigo, Ind. Eng. Chem. Res. 2002, 41 , 4187–4194. 28. P. Canizares, J. Garcia-Gomez, J. Lobato, M.A. Rodrigo, Ind. Eng. Chem. Res. 2004, 43 , 1915–1922. 29. P. Canizares, J. Garcia-Gomez, J. Lobato, M.A. Rodrigo, Ind. Eng. Chem. Res. 2004, 43 , 1923–1931. 30. O. Scialdone, A. Galia, C. Guarisco, S. Randazzo, G. Filardo, Electrochim. Acta 2008, 53 , 2095–2108. 31. O. Scialdone, A. Galia, S. Randazzo, in preparation. 32. M. Bergmann, J. Rollin, Catalysis Today 2007, 124 , 198–203. 33. L. Szpyrkowicz, J. Naumaczky, F. Zilio-Grandi, Toxicolog. and Environ. Chem. 1994, 44 , 189–202. 34. C. Comninellis, A. Nerini, J. Appl. Electrochem. 1995, 25 , 23–28. 35. L.C. Chiang, J.E. Chang, T.C. Wen, Water Res. 1995, 29 , 671–678. 36. C.H. Yang, C.C. Lee, T.C. Wen, J. Appl. Electrochem. 2000, 30 , 1043–1051. 37. A.M. Polcaro, A. Vacca, M. Mascia, F. Ferrara, J. Appl. Electrochem. 2008, 38 , 979–984; A.M. Polcaro, A. Vacca, M. Mascia, S. Palmas, J.R. Ruiz, J. Appl. Electrochem. 2009, 39 , 2083–2092. 38. F. Bonfatti, S. Ferro, F. Lavezzo, M. Malacarne, G. Lodi, A. De Battisti, J. Electrochem. Soc. 2000, 147 , 592–596. 39. C.A. Martinez-Huitle, S. Ferro, A. De Battisti, Electrochem. Solid State Lett . 2005, 11 , D35–D39. 40. O. Scialdone, S. Randazzo, A. Galia, G. Silvestri, Water Res. 2009, 43 , 2260–2272. 41. K. Serrano, P.A. Michaud, C. Comninellis, A. Savall, Electrochim. Acta 2002, 48 , 431–436. 42. O. Scialdone, S. Randazzo, A. Galia, G. Filardo, Electrochim. Acta 2009, 54 , 1210–1217. 43. M. Wu, G. Zhao, M. Li, L. Liu, D. Li, J. Hazardous Mat . 2009, 163 , 26–31. 44. L. Szpyrkowicz, M. Radaelli, J. Appl. Electrochem. 2006, 36 , 1151–1156. 45. A. Anglada, A. Urtiaga, I. Ortiz, Environ. Sci. Technol. 2009, 43 , 2035–2040. 46. A. Anglada, A. Urtiaga, I. Ortiz, J. Chem. Technol. Biotechnol . 2009, 84 , 1747–1755. 47. M. Zhou, H. S¨arkk¨a, M. Sillanp¨aa¨ , Separ. Purif. Techn. 2011, 78 , 290–297.

12 Production of Strong Oxidizing Substances with BDD Anodes Ana S´anchez-Carretero, Cristina S´aez, Pablo Canizares, and ˜ Manuel A. Rodrigo

12.1

ELECTROLYSES WITH CONDUCTIVE-DIAMOND ANODES

From the late nineties of the twentieth century to present, applications of conductive diamond surfaces in electrochemistry have grown significantly. One of these applications is their use as anodes in electrochemical wastewater treatment processes, where these anodes have led to very powerful oxidation processes, with high efficiencies in the removal of organic pollution. During the characterization of these processes, there have been reports about the significant action of mediated electro-oxidation in the achievement of good results, and the role of many particular oxidants has been described. In particular, the role of hydroxyl radical has been established to be very relevant. Its occurrence during diamond electrolyses of aqueous wastes was demonstrated by Marselli et al. [1], and later works have confirmed this occurrence with very different observations. This allowed classifying conductivediamond electrochemical oxidation (CDEO) of wastewaters as an advanced oxidation process (AOP). In this context, one of the more illustrative observations [2,3] about the role of hydroxyl radicals in wastewater treatment compares results of the electrolyses of phenol working at anodic potentials below and over the necessary amounts for water oxidation (around 2.5–2.7 V versus NHE). In this way, Figure 12.1 shows a comparison of the electrochemical windows of Pt and BDD electrodes and also summarizes the anodic Synthetic Diamond Films: Preparation, Electrochemistry, Characterization, and Applications, First Edition. Edited by Enric Brillas and Carlos Alberto Mart´ınez-Huitle. © 2011 John Wiley & Sons, Inc. Published 2011 by John Wiley & Sons, Inc.

281

282

PRODUCTION OF STRONG OXIDIZING SUBSTANCES WITH BDD ANODES

Current density(A m–2)

Anodic potential (V) vs. SCE

20

2.0

65

2.2

90

2.5

155

3.5

195

4.1

0.04 Thermodynamics

0.03

Pt

j (A cm–2)

0.02

DDB

0.01 0

–0.01 –0.02

–2.0

–1.0

0.0

1.0

2.0

3.0

E (V) vs. NHE

–0.03 Figure 12.1 Experimental methodology to demonstrate the role of hydroxyl radicals in CDEO process. (Adapted from Ref. 3.) See color insert.

potential measurement during the conductive-diamond phenol oxidation using different current densities. Figure 12.2 shows how the electrolyses of phenol solutions at potentials below water electrolyses (and hence without formation of hydroxyl radicals) lead to the formation of significant amount of intermediates (aromatic and aliphatic). Electrolyses at potentials higher than this value lead to the formation of negligible intermediates and to high efficient processes, suggesting that hydroxyl radicals improve the oxidation conditions and lead to very strong oxidation conditions. However, the better performance of conductive-diamond electrolyses over other advanced oxidation processes (those based on the production and use of hydroxyl radicals), and also over other electrolytic processes, indicates that in addition to direct oxidation and to hydroxyl radicals-mediated oxidation, there should be many other oxidants with a significant role in the destruction of organics during electrochemical wastewater oxidation with diamond anodes. In this context, the roles of chlorine [4–7], sulphates [8–10], phosphates [11,12], and many other types of salts on the electrochemical destruction of organics have been extensively studied in the literature. According to this fact, Figure 12.3 summarizes the main mechanism proposed to explain the good results of CDEO in the oxidation of organic pollutants during wastewater treatment. From this figure, it seems clear that the electrolyses of water solutions of different salts can lead to the formation of oxidants, and if proper conditions are found, oxidants produced can be separated and stored for later use (see Figure 12.4). This presence of oxidant species has encouraged many research groups to study the synthesis of particulate oxidants and to isolate them as valuable products. This work is intended to review some of these processes as examples of the significant applications of CDEO in the production of high-value oxidants. To do this, the

350

1200

300

1000

250

COD/g m–3

TOC(g m–3)

12.2 PRODUCTION AND STORAGE OF OXIDIZING SUBSTANCES: EXPERIMENTAL SETUPS

200 150 100

800 600 400 200

50

0

0 0

500

1000

1500

0

2000

Time(min)

500

1000 1500 Time(min)

2000

350 Carboxylic acids/g Cm–3

70 Aromatics(g Cm–3)

283

60 50 40 30 20 10

300 250 200 150 100 50 0

0 0

10

30 20 Q (kAh m–3)

40

50

0

20 40 Q(kAh m–3)

60

Figure 12.2 Results of the electrolyses of phenol solutions at different anodic potentials (C0 : 6.4 mM phenol; pH 7; T:25◦ C). (+) 2.0 V versus NHE, (◦) 2.2 V versus NHE, (△) 2.5 V versus NHE, () 3.5 versus NHE, () 4.1 V versus NHE. (Adapted from Ref. 2.)

production of hydroxyl radicals and more stable oxidants (peroxosalts and peroxoacids, ferrates, and halogen oxoanions) are going to be described. 12.2 PRODUCTION AND STORAGE OF OXIDIZING SUBSTANCES: EXPERIMENTAL SETUPS Prior to the description of processes that produce oxidants, it is necessary to take into account that contrary to the electro-oxidation of wastewaters, the synthesis of oxidants should be carried out in a double-compartment electrochemical flow cell in which anodic and cathodic compartments are separated by means of an ionic exchange membrane. This avoids the cathodic reduction of oxidants, and hence increases significantly the efficiency of the process. Usually, the ionic membrane is a cathionic one, because many oxidants are anionic species, and this allows storing the produced oxidants in the anodic compartment. Figure 12.5 illustrates with an example (obtained during the electrolytic production of perphosphates with boron-doped diamond anodes) the improvement obtained in the production of oxidants with the use of double compartment cells. At this point, it is important to be reminded that the use of double compartment cells in the destruction of the organic pollutants contained in wastewaters is not necessary because most reactions are irreversible. The separation of the anionic and cationic compartments only leads to an increase in the cell potential, and consequently in the operation cost, without any clear improvement in the treatment results.

284

PRODUCTION OF STRONG OXIDIZING SUBSTANCES WITH BDD ANODES

HYDROXYL RADICALS MEDIATED ELECTROLYSES

Anode DIRECT ELECTROLYSES

Direct

Si

Si e–

Si+1

With oxidants produced from saltsas mediators

Si+1

Mred-Mox

Si

Mred

2SO42– → S2O82–

H2O e–

Si+1

Mox

2PO43– → P2O84– Cl– → ClO–

OH· With ozone as mediator

O2

Si

2H+ + O2 Oxygen evolution

O3

Mred

H2O

Si

e–

Mox

Si+1

H2O2

Si+1 With hydrogen Si peroxide as mediator

Si+1

Mediated electrolyses with oxidants produced from salts contained in the wastewater Figure 12.3 Main mechanisms for the electrolyses with conductive-diamond anodes of wastewaters polluted with organics.

Another important point to be considered in the production of oxidants is their stability. Light, temperature, and pH use have a very significant influence on this stability, which is usually favored at low temperature, dark, and extreme pHs. An example of the huge influence of these parameters in the stability of oxidants is shown in Figure 12.6, in which the decomposition rate of peroxophosphates is shown as a function of the pH and temperature. For this reason, a typical setup to produce oxidants with diamond anodes should include a cryostat device to control temperature, a pH control device, and dark glass in tubes and storage tanks. With such requirements, an example of a complete setup to produce oxidants with conductive-diamond anodes electrolyses is shown in Figure 12.7. 12.3 PRODUCTION OF HYDROXYL RADICALS WITH CONDUCTIVE-DIAMOND ANODES Production of hydroxyl radicals during conductive-diamond electrolysis of aqueous wastes was first demonstrated by Marselli and co-workers [1] using selective reactions of hydroxyl radicals. This was a very significant work because it was the first

12.3 PRODUCTION OF HYDROXYL RADICALS WITH CONDUCTIVE-DIAMOND ANODES

285

Anode

Mred H2O

Mred-Mox

e–

Mox 2SO42– → S2O82–

OH·

2PO43– → P2O84–

O2 2H+ + O2

Cl– → ClO–

Oxygen evolution

O3 H2 O

Mred e–

H2O2

Mox

Figure 12.4 anodes.

Mechanisms for the production of oxidants with electrolyses with conductive-diamond

Oxidants(mmol) P2O84–

300

Double-compartment electrochemical flow cell

250 200 150 100

Single-compartment electrochemical flow cell

50 0 0

50

100

150

200

250

300

Time(min)

Figure 12.5 Effect of cell compartments on the efficiency of the production of oxidants by electrolysis with conductive diamond anodes. (C0 : 1 M K3 PO4 ; pH 12.5; T: 25◦ C; j: 1250 A m−2 ).

286

PRODUCTION OF STRONG OXIDIZING SUBSTANCES WITH BDD ANODES

pH Decomposition rate(mol (1 min)–1)

1.00E–05 0

1

2

3

4

5

6

7

1.00E–06

1.00E–07 Figure 12.6 Effect of temperature and pH on the decomposition rate of peroxophosphates. (C0 : 1 M HK3 PO4 ; j: 64 A m−2 ). (+) 35◦ C, (△) 22◦ C, () 10◦ C. (Reprinted from Ref. 36.)

Power supply V

A

catholyte

anolyte +

–

Electrochemical cell

Membrane Anode (BDD)

Out

In

Cathode (AISI 304)

Out

In

Figure 12.7 Typical setup to produce oxidants with CDEO technology. See color insert.

12.3 PRODUCTION OF HYDROXYL RADICALS WITH CONDUCTIVE-DIAMOND ANODES

287

demonstration that conductive-diamond electrochemical oxidation belongs to the group of the advanced oxidation technologies. Figure 12.8 summarizes graphically the two main findings of the work: the selective formation of DMPO and the selective hydroxylation of salicylic acid. The selective oxidation of 5,5-dimethyl-1-pyrroline-N-oxide (DMPO) involves trapping the hydroxyl radicals by an addition reaction (spin trapping) to produce a more stable radical spin adduct. The anodic oxidation of salicylic acid at the BDD anode leads to the formation of two particular dihydroxylated products, 2,3- and 2,5-dihydroxybenzoic acids, just the same products formed by Fenton reaction. After that, many other works have supported this achievement with different observations that indicate the significance of hydroxyl radicals in the electrolyses with conductive diamond. The main drawback of hydroxyl radical is their small stability, which makes its storage impossible and even makes their detection very difficult (average lifetime lower than 200 µs [13]). However, the occurrence of hydroxyl radicals explains the production of many more stable oxidants, particularly two relevant oxidants: ozone and hydrogen peroxide. Thus, although many other technologies can provide for ozone or hydrogen peroxide in a more efficient way, there are many works in literature in which the production of both oxidants is obtained with conductive-diamond electrolyses. As it has been describe previously, hydroxyl radicals are produced during conductivediamond electrolyses of aqueous solutions. These radicals decompose following a complex system of reactions in which ozone [14] and hydrogen peroxide [14] can be intermediates or final products. In literature, it is reported that in absence of organic compounds, hydroxyl radicals can react with each other to form hydrogen peroxide from Reaction (12.1) [15]. In fact, Michaud et al. [16] observed the generation of hydrogen peroxide during the electrolysis with BDD anodes of HClO4 solutions. In addition, it is reported that electro-generated hydrogen peroxide can be further oxidized to oxygen either by its direct discharge on the

Figure 12.8 Demonstration of the occurrence of hydroxyl radicals during CDEO.

288

PRODUCTION OF STRONG OXIDIZING SUBSTANCES WITH BDD ANODES

electrode surface by Reaction (12.2) or by hydroxyl radical-mediated Reaction (12.3). •

OH · + • OH· → H2 O2 +

H2 O2 → O2 + 2H + 2e

(12.1) −

H2 O2 + • OH· → O2 + 2H2 O

12.4

(12.2) (12.3)

SYNTHESIS OF PEROXOACIDS AND PEROXOSALTS

Peroxoacids and peroxosalts are in a group of oxidants characterized by the presence of a group –O–O– in the molecules. Typically, these species come by the substitution of a group –OH by a –O–OH in oxoacids or oxosalts of the groups IV, V, and VI of the periodic table. Typical examples include peroxocarbonates, peroxonitrates, peroxosulphates, and peroxophosphates, although the last two are more important because of their actual and potential applications. The efficiency in the production of these products is high enough to study commercial applications; consequently, many applications have been proposed for these species. The main use of peroxosulphuric acids is as an initiator in polymerization processes (olefins, vinyl chloride, styrene-butadiene, vinyl acetate, acrylic ester). Its use as a reagent in the etching of printed circuit boards and in the removal of photo resists are also important. Other applications concern dyes oxidation, whitening fibers, promotion for radical polymerization, total organic compound measurements, and so on. Peroxophosphates have a wide variety of applications in areas such as oxidizing agents in organic synthesis [17,18] cosmetics [19,20], agriculture [21], wastewater treatment [22], and also as bleaching agents in the detergent industry [23,24]. Due to its similar properties, peroxodiphosphate can also be used in other processes as a substitute of peroxodisulphate. In these later uses, peroxodiphosphate has an important advantage: It is a more environmental-friendly reagent because its reduction product (phosphate) can be easily and economically removed from aqueous wastes that the reduction product of persulphates (sulphates). Electrolyses with conductive-diamond anodes have shown to be effective in the production of these species. In the following subsections, some details about the electrolytic production with diamond anodes of peroxodisulphuric acid, monoperoxophosphoric acid, and peroxodiphosphate are going to be described. 12.4.1

Peroxosulphuric Acids

Peroxosulphuric acids can be produced by electrolytic oxidation of sulphuric acid solutions. Two different species are included in this group: peroxomonosulphuric and peroxodisulphuric acid (Figure 12.9). They both have a very high reduction potential (1.81 and 2.08 V, respectively). Peroxodisulphuric acid is more stable than peroxomonosulphuric acid, but it can also suffer thermal decomposition giving to the generation peroxomonosulphuric acid and also hydrogen peroxide. The process is more significant at temperatures higher than 50◦ C, but it can be observed at any range of temperatures, and hence mixtures of both acids are typically encountered in commercial products in which the dimmer is the primary species.

12.4 SYNTHESIS OF PEROXOACIDS AND PEROXOSALTS

O HO

S O OH

O Monoperoxosulphuric acid

O

289

O

HO S O O S OH O O Peroxodisulphuric acid

Figure 12.9 Peroxosulphuric acids.

Traditionally, the synthesis of peroxosulphuric acids have been carried out by electrolyses of sulphuric acid by Reaction (12.4) with noble metal-coated anodes such as platinized, titanium, tantalum, or niobium [25]. 2H2 SO4 → H2 S2 O8 + 2e−

(12.4)

Efficiency of the process was found to depend strongly on the electrode material selected, and particularly on the high overvoltage value for oxygen evolution reaction provided by the anode material. In this context, the use of conductive-diamond electrodes was first proposed by Gandini et al. [26] and by Michaud et al. [27] in 2000, being one of the first electrochemical applications proposed for the anodes of conductive diamond. Figure 12.10 shows the results of the electrolyses with diamond electrodes of sulphuric acid solutions. As it can be observed, the production of significant concentrations of peroxosulphuric acids is obtained, and the efficiency is high enough to be used commercially. During the batch electrosynthesis, the oxidant concentration increases with the specific charge until a constant value is achieved, and current efficiency decreases continuously. This behavior could be explained in terms of mass transfer limitations (in the batch system studied, the concentration of reactant decreases continuously) or, most likely, some sort of chemical or electrochemical destruction of the oxidants formed, which promotes the concentration of a formed product. Electrolytic production of peroxosulphuric acids was found to depend significantly on the concentration of raw sulphuric acid, temperature, and current density. The effect of these parameters is shown in Figures 12.11 and 12.12, respectively. The production of peroxosulphuric acids and the efficiency of the processes increase with the concentration of raw H2 SO4 , with low temperatures and with high current densities. The influence of concentration is difficult to explain although it is as expected: The higher the concentration of raw materials, the higher the concentration of product. Preliminarily, it could be explained by mass transfer controlling mechanisms. However, concentration of sulphuric acid does not become small in any case, and hence some sort of electrochemical decomposition of peroxosulphuric acid should be considered to explain the effect of concentration. This also has to be related to the stabilization of the concentration of peroxosulphuric acids during batch electrolysis, and the plateau range of current charge could indicate the zone in which the production rate is equal to the decomposition rate. The influence of temperature is easier to explain. It is related with the thermal stability of peroxosulphuric acids, which are known to decompose with temperature to yield

290

PRODUCTION OF STRONG OXIDIZING SUBSTANCES WITH BDD ANODES

Oxidants(mmol H2S2O8)

40

30

20

10

0 0

20

40

60

80

100

Q (Ah dm–3)

Current efficiency(%)

40

30

20

10

0 0

20

40 60 Q (Ah dm–3)

80

100

Figure 12.10 Variation of the oxidants () and of current efficicency () with the specific electrical charge passed in the electrosynthesis of peroxosulphuric acids (C0 : 2 M H2 SO4 ; T: 10◦ C; j: 1250 A m−2 ).

sulphuric acid and hydrogen peroxide according to the mechanisms proposed in Reactions (12.5) to (12.7) [28]. S2 O8 2− + H2 O → 2 SO4 2− + 2H+ + 1/2O2

(12.5)

2− + S2 O8 2− + H2 O → SO2− 5 + SO4 + 2H

(12.6)

SO5 2− + H2 O → H2 O2 + SO4 2−

(12.7)

The influence of current density is the more interesting fact in the production of peroxosulphuric acids: There are two different behaviors as a function of the current density. Current densities higher than 1100 A m−2 lead to a more efficient process. Initially, this observation could be related to the massive formation of hydroxyl radicals in the reaction media, which complements the direct electrochemical formation of peroxosulphuric acid

12.4 SYNTHESIS OF PEROXOACIDS AND PEROXOSALTS

291

Oxidants(mmol H2S2O8 dm–3)

30 25 20 15 10 5 0

0

0.5 1 1.5 Concentration raw H2SO4(M)

2

Figure 12.11 Effect of the concentration of raw sulphuric acid on the production of peroxosulphuric acids for a specific current charge of 40 Ah dm−3 . ( j: 1250 A m−2 , T: 10◦ C).

Oxidants(mmol H2S2O8 dm–3)

30 25 20 15 10 5 0 0

10

20 30 Temperature(°C)

40

50

Oxidants(mmol H2S2O8 dm–3)

30 25 20 15 10 5 0 0

500

1000 1500 Current density(A m–2)

2000

Figure 12.12 Effect of the operation conditions on the production of peroxosulphuric acids for a current charge of 40 A h dm−3 . (C0 : 1 M H2 SO4 , pH 2).

292

PRODUCTION OF STRONG OXIDIZING SUBSTANCES WITH BDD ANODES

from Reactions (12.8) to (12.10). As it is going to state with other oxidants, current densities around 1100 A m−2 correspond to anodic potentials for water electrolysis in the experimental setup used. HSO4 − + • OH → SO4 − • + H2 O SO4

2−

•

+ OH → SO4

−•

+ H3 O

(12.8)

+

(12.9)

SO4 − • + SO4 − • → S2 O8 2− 12.4.2

(12.10)

Peroxodiphosphate Salts

Peroxophosphates chemistry is similar than that of persulphates, although peroxophosphates are less known because their synthesis methods were not very efficient, and they usually led to unpure species, but contaminated with some of the additives dosed to increase the efficiency of the production process. Similar to peroxosulphuric acid chemistry, two different species can be found: peroxomonophosphate and peroxodiphosphate salts (see Figure 12.13). Peroxomonophosphate is stable at acid pH, whereas peroxodiphosphate stability is higher at basic pH. Both are powerful oxidants, being their oxidation potentials are similar to those of persulphates (2.07 V peroxodiphosphate versus 2.08 V persulphates). Typically, potassium peroxodiphosphate, K4 P2 O8 , is obtained by electrolysis on platinum electrodes at alkaline conditions of potassium phosphate solutions containing some reagents [29], mainly fluoride or thiocianate. These reagents are primarily added to promote the blockage of oxygen evolution sites [30] of the anode in the synthesis process. Consequently, the direct oxidation of phosphate to peroxodiphosphate is favored over the water oxidation process and higher current efficiencies are obtained. However, some of the uses of peroxodiphosphates can be affected by the impurities. Likewise, some of the reagents can be highly corrosive to platinum and others can be toxic. Thus, in order to obtain impurities-free peroxodiphosphate, extensive purification is required [31]. This greatly increases the manufacturing costs. The synthesis of K4 P2 O8 can also be carried out with platinum without using any additives, but in these cases low current efficiencies are obtained [32]. Figure 12.14 shows the production of peroxodiphosphate by electrolysis with conductive-diamond of K3 PO4 in alkaline conditions (pH 12.5; T: 25◦ C; j: 1250 A m−2 ) [33]. As it can be observed, the process results are surprising. Huge conversions are obtained with very good efficiencies. In addition, peroxodiphosphate can be precipitated and separated from the reaction media simply by the addition of methanol to the electrolyzed solution. It can also be observed that the rate of generation of peroxodiphosphate decreases progressively during a batch process. This behavior (just the same that was observed in the electrochemical production of peroxosulphuric acids OH O P O OH OH Monoperoxophosphoric acid

OK

OK

O P O O P O OK

OK

Potassium peroxodiphosphate

Figure 12.13 Peroxophosphoric acids and salts.

12.4 SYNTHESIS OF PEROXOACIDS AND PEROXOSALTS

293

100

Conversion(%)

80 60 40 20 0 0

30

60

90

120

150

180

210

Q (Ah dm–3)

Current efficiencies(%)

100

80

60

40 20

0 0

30

60

90

120

150

180

210

Q (Ah dm–3) Figure 12.14 Variation of the conversion and of the current efficiency with the specific electrical charge passed in the electrosynthesis of peroxodiphosphate (C0 : 1 M K3 PO4 ; pH 12.5; T: 25◦ C; j: 1250 A m−2 ). (Reprinted from Ref. 33.)

with diamond anodes) is justified in terms of a progressive decrease in the efficiency of the process. This continuous decrease can be due to mass transfer limitations or some sort of electrochemical destruction of the oxidants formed. In this case, the high conversions obtained and the high stability of peroxodiphosphate suggests that mass transfer limits the process. Figure 12.15 shows the influence of the raw material used to produce peroxodiphosphate, particularly the concentration of potassium phosphate and the pH. As expected, process efficiencies increase with the concentration of raw phosphate in the solution to be electrolyzed and hence for the same current charge passed the production of peroxodiphosphate increases with the initial concentration of phosphates. However, on a few occasions it was observed that electrolyses of solutions with very high concentrations of K3 PO4 seemed to damage to the diamond surface and small corrosion circles appeared on the surface of the diamond after the treatment. This problem never appeared

294

PRODUCTION OF STRONG OXIDIZING SUBSTANCES WITH BDD ANODES

90 80

Conversion(%)

70 60 50 40 30 20 10 0 0

1 Concentration of raw potassium phosphate/M

2

90 80 Conversion(%)

70 60 50 40 30 20 10 0 8

9

10

11 pH

12

13

14

Figure 12.15 Effect of the characteristic of the raw materials on the electrosynthesis of peroxodiphosphate at an electrical charge passed of 130 Ah dm−3 (T: 25◦ C; j: 1250 A m−2 ). (Reprinted from Ref. 33.)

when working with initial concentration of 1 M K3 PO4 ; hence, this concentration should be recommended to warrant good results, and to protect the anodic surface. Likewise, pH influences greatly in the process performance and only good yields are obtained in strongly alkaline solutions. Optimum operation pH was 12.5—the same that literature indicates is the maximum stability of peroxodiphosphate. Figure 12.16 shows the effect of the main operation parameters (temperature and current density) on the efficiency of the process. Temperatures higher than 25◦ C lead to low conversions, and thus, to low current efficiencies. The thermal decomposition of peroxodiphosphate to give pyrophosphate and oxygen from Reaction (12.11) might justify the observed decrease, although the complex chemistry of the system needs more chemical studies to clarify this point. P2 O8 4− → P2 O7 4− + 1/2O2

(12.11)

12.4 SYNTHESIS OF PEROXOACIDS AND PEROXOSALTS

295

80 70

Conversion(%)

60 50 40 30 20 10 0 10

15

20

25 30 Temperature(°C)

35

40

45

100

Conversion(%)

80 60 40 20 0 0

50

100

150

200

250

j(mA cm–2) Figure 12.16 Effect of the operation conditions on the electrosynthesis of peroxodiphosphate with diamond anodes at an electrical charge passed of 80 Ah dm−3 (C0 : 1 M K3 PO4 ; pH 12.5). (Reprinted from Refs. 33,41.)

The influence of the current density is the same as that observed for peroxosulphuric acids. It can be seen that the conversion increases continuously with the specific charge passed, and that for a given current charge passed not a continuous change is obtained as a function of the current density, but only two limit behaviors can be discerned. For current densities below 1000 A m−2 maximum conversions do not exceed 30%, whereas for higher current densities, conversions over 70% are obtained. The abrupt change in the efficiency may be related to the mechanisms of peroxodiphosphate formation on BDD surfaces (see Figure 12.17). Thus, it is proposed [30] that peroxodiphosphate can be formed by direct electrooxidation from Reaction (12.12) on the surface of electrodes such as platinum. 2PO4 3− → P2 O8 4− + 2e−

(12.12)

296

PRODUCTION OF STRONG OXIDIZING SUBSTANCES WITH BDD ANODES

e–

PO4–3

anode

(PO4–2)• (PO4–2)• e–

H2O OH

Direct oxidation P2O8

–4

PO4–3 (PO4–2)•

(PO4–2)•

P2O8–4

OH mediated oxidation

Figure 12.17 Mechanism for the production of peroxodiphosphate with diamond anodes.

However, the higher efficiencies obtained with BDD anodes can be due to the presence of hydroxyl radicals. It is reported [1] that large quantities of these radicals are formed during electrolyses of aqueous solutions. These radicals can combine with phosphate ions and form PO4 2− • radicals [30]. These later radicals can oxidize other compounds in a region close to the anode surface (e.g., water peroxide to oxygen), or they can combine between them to form peroxodiphosphate or with hydroxyl radicals to form peroxomonophosphate. This can justify the higher efficiencies obtained in the electrosynthesis of these compounds when using BDD as anode materials, as these anodes combine both direct and hydroxyl-mediated oxidation processes. Conversely, this complex mechanism might justify the abrupt change in the efficiencies by the promotion of one or the two mechanisms in the electrochemical oxidation. Nevertheless, further research should be done in order to clarify this point. 12.4.3

Monoperoxophosphoric Acid

The synthesis of peroxomonophosphoric acid (H3 PO5 ) is usually based on highly exothermic chemical reactions, which, as a major drawback, led to the rapid decomposition of the peroxo-compounds generated and consequently to a low efficiency in the production of monoperoxophosphoric acids. Several improvements, such as the addition of inert diluents, have been proposed [34], but the low reproducibility of the results limits their use. For this reason, the hydrolysis of peroxodiphosphate salts in very acidic conditions appeared as one of the best synthesis method of peroxomonophosphoric acid [35]. However, the requirement of powerful reagents (peroxodiphosphate and perchloric acid or hydrogen peroxide), and its large manufacturing costs have limited its commercial application. In this context, the direct electrochemical production of monoperoxophosphoric acid using conductive-diamond electrodes can become in an important way of manufacturing this product. Figure 12.18 shows the results of batch electrolyses with conductive diamond of phosphoric acid [36]. As shown, monoperoxophosphoric acid is formed. As in the synthesis of other compounds, oxidant concentration increases to achieve a plateau. This plateau value results from the compensation of the rate of formation of

12.4 SYNTHESIS OF PEROXOACIDS AND PEROXOSALTS

297

4.5

Oxidants(mmol PO53–)

4 3.5 3 2.5 2 1.5 1 0.5 0 0

20

0

20

40 Q (Ah dm–3)

60

80

4.5 4

Efficiency(%)

3.5 3 2.5 2 1.5 1 0.5 0 40 Q (Ah dm–3)

60

80

Figure 12.18 Variation of the monoperoxophosphoric acid concentration and current efficiencies with electrical charge passed in the electrolysis of phosphoric acid (C0 : 1 M H3 PO4 ; pH 1.2; T: 13◦ C; j: 64 A m−2 ). (Reprinted from Ref. 36.)

oxidant and its decomposition rate. The results obtained exceed those obtained with other methods. However, the efficiency and the amount of oxidant generated are significantly smaller than the production of peroxodiphosphates salts using electrolyses of alkaline potassium phosphate. In spite of that, the efficiencies are high enough to suggest that this method could be a promising alternative to acidification of peroxodiphosphate solutions. Figure 12.19 shows the influence of the raw phosphoric acid (concentration and pH of the solution) on the production of monoperoxophosphoric acid for a particular electrical current charge passed. With respect to the H3 PO4 concentration, the amount of oxidant obtained is largely influenced by the concentration of the solute. The quantity of peroxomonophosphate increases significantly with the amount of H2 PO4 − ions available to be oxidized until achieving a given value (around 1 M H3 PO4 ). and then a decrease in the amount of peroxomonophosphate produced is observed. The optimum concentration for the peroxomonophosphate synthesis with BDD anodes in the operating conditions used in this work is 1 M H3 PO4 .

298

PRODUCTION OF STRONG OXIDIZING SUBSTANCES WITH BDD ANODES

Oxidants(mmol PO53–)

1 0.8 0.6 0.4 0.2 0 0

0.5

1

1.5

2

2.5

3

5

6

Initial concentration(mol dm–3)

Oxidants(mmol PO53–)

5 4 3 2 1 0 0

1

2

3

4

pH Figure 12.19 Amount of oxidant generated in the electrosynthesis of peroxomonophosphate as a function of the initial concentration of H3 PO4 at an electrical charge passed of 1.5 Ah dm−3 (pH 5; T: 13◦ C; j: 20 A m−2 ). (Reprinted from Ref. 36.)

This unexpected behavior has to be explained in terms of the hydroxyl radical’s role in the formation of monoperoxophosphoric acid. Thus, hydroxyl radicals can contribute to the generation of H2 PO4• . This radical can be produced by direct electrolyses from Reaction (12.13) on the electrode surface or by the action of hydroxyl radicals from Reaction (12.14) in the nearest of the electrode. Moreover, the presence of hydroxyl radicals is also needed to promote the generation of peroxomonophosphoric acid by Reaction (12.15). As a result, an increment of the initial concentration of phosphate is not enough to ensure the generation of higher amount of oxidant generated, but the coexistence of both radical reagents is required. H2 PO4 − → (H2 PO4 ) • + e−

(12.13)

H2 PO4 − + • OH· → (H2 PO4 ) • + OH− •

•

(12.14) 9

−1 −1

(H2 PO4 ) + OH → H3 PO5 k = 4 · 10 L mol s

(12.15)

12.4 SYNTHESIS OF PEROXOACIDS AND PEROXOSALTS

299

Conversely, the pH seems not to have a strong influence on the results, and within the range of 1 to 5 almost no changes are observed. This is important because due to the acid-base equilibrium given by Reactions (12.16) to (12.18), the hydrogenphosphoric anion can also be precursor of the formation of monoperoxophosphoric acid by Reaction (12.19). H3 PO4 ↔ H2 PO4 − + H+

pKa = 2.14

(12.16)

H2 PO4 − ↔ HPO4 2− + H+

pKa = 6.86

(12.17)

HPO4 2− ↔ PO4 3− + H+

pKa = 12.4

(12.18)

HPO4 2− → (HPO4 − ) • + e−

(12.19)

Figure 12.20 shows the influence of the operation conditions (temperature and current density) on the process at particular current charge passed. As it can be observed, the

Oxidants(mmol PO53–)

4

3

2

1

0 0

5

10

15 20 Temperature(°C)

25

30

35

Oxidants(mmol PO53–)

5 4 3 2 1 0 0

300

600

900

1200

1500

Current density(Am−2)

Figure 12.20 Amount of oxidant generated in the electrosynthesis of peroxomonophosphate as a function of the temperature at an electrical charge passed of 10 Ah dm−3 (1 M H3 PO4 ; pH 5; j: 64 A m−2 ). (Reprinted from Refs. 36,41.)

300

PRODUCTION OF STRONG OXIDIZING SUBSTANCES WITH BDD ANODES

conversion of phosphoric acid to monoperoxophosphoric acid shows a marked dependence with this operating parameter. The higher the temperature, the lower the oxidant concentrations obtained at the steady state. The influence of the current density is just as reported for the oxidants previously described in this chapter. Two significantly different types of behavior can be discerned. At low current densities, a smaller efficiency suggests that the direct mechanism is the main responsible for peroxoacid generation. At higher intensities, the mechanism may involve also the hydroxyl radicals as it was explained previously.

12.5 12.5.1

SYNTHESIS OF HALOGEN OXOANIONS Perchlorates

The good properties in the production of oxidants have encouraged the interest of comparing its performance with the performance of dimensionally stable anodes (DSA) in the production of hypochlorite. However, the properties of conductive diamond are completely different from that of DSA. Thus, Figure 12.21 shows the variation of the concentration of the main oxidant species generated during the electrolysis of 0.1 M NaCl solutions with DSA and BDD anodes. As it can be observed, hypochlorite, chlorite, and chlorate are the species produced in higher concentrations for low current charges in the case of conductive-diamond anodes [37–39]. Likewise, from a given electrical charge passed, significant amounts of perchlorate begin to be detected in the reaction system and the conversion to this product is quantitative for large current charges. These results are opposite to those obtained with other electrodes such as DSA in which chloride ions are almost only transformed into hypochlorite ions. It is well documented [40] that DSA electrodes can be successfully used for the electrochemical generation of hypochlorite, although they are not able to attain significant concentrations of perchlorate salts. 12.5.2

Perbromates

Perbromates are oxoanions of bromine that cannot be presently found commercially due to the nonexistence of a good synthesis method. Figure 12.22 shows results about the production of perbromate during the oxidation of bromate solutions at strong alkaline media. It can be easily verified that conductive diamond allows the further oxidation of bromate ions, giving to the formation of Br(VII) species [41]. This observation is a very important insight because, to the authors’ knowledge, the electrochemical generation of perbromate has not been previously reported by electrochemical methods in aqueous media. However, at the present moment, the low concentration attained and the low current efficiency serve as warnings about the use of this electrochemical technique as a promising synthesis method. So, further studies must be done to establish the optimum operating conditions that allow improving the current efficiency and the generation of large amount of perbromate.

12.6 SYNTHESIS OF FERRATES

301

90 DSA 80

% Conversion

70 60 50 40 30 20 10 0

0

50

100

150

Q (Ah dm–3) 40

BDD

35

% Conversion

30 25 20 15 10 5 0 0

5

10

15

20

25

30

Q (Ah dm–3) Figure 12.21 Production of perchlorates with dimensionally stable anodes (DSA) and conductive− diamond anodes (DDB) (C0 : 0.1 M NaCl, pH 10, T: 35◦ C, j: 300 Am−2 ). () ClO− , (•) ClO− 3 , (△) ClO2 , . () ClO− 4

12.6

SYNTHESIS OF FERRATES

Typical oxidation states for iron species are +2 and +3; however, unusual oxidation states are observed to be stable in particular conditions, such as +6 in the form of ferrate (VI), FeO4 2− . Ferrate ions are very powerful oxidizing agents with standard halfcell reduction potentials ranging from 2.20 V at acidic pHs to 0.72 V versus NHE at alkaline condition [42,43]. Moreover, during the oxidation process, ferrate (VI) ions will be reduced to Fe (III) ions or ferric hydroxide, making them suitable to be used in a wide

302

PRODUCTION OF STRONG OXIDIZING SUBSTANCES WITH BDD ANODES

12

Oxidants(mmol BrO4–)

10 8 6 4 2 0 0

100

200

300

400

Q (Ah dm–3) Figure 12.22 Production of perbromates with conductive diamond anodes. (C0 : 0.1 M NaBrO3 , pH 10, T: 17◦ C, j: 1000 Am−2 ).

range of applications: “green” organic synthesis [44,45], wastewater treatment (oxidation and coagulation) [46,47], and water treatment (persistent disinfection, coagulation, and oxidation of nondesirable compounds) [42,48–50]. Due to their highly oxidized iron basis, multiple electron transfer, and high intrinsic energy, ferrate (VI) can be used as “super-iron” catode [45]. The ferrate synthesis can be divided into the following three categories: (1) thermal chemical synthesis, by heating/melting various iron oxide containing minerals under conditions of strong alkaline and oxygen flow [51]; (2) wet chemical synthesis, by oxidizing Fe(III) salt at strong alkaline condition and using hypochlorite or chlorine as the oxidant [52] or electrochemical techniques; and (3) anodic oxidation using iron or alloy as anode and NaOH or KOH as electrolyte [53]. However, these generation techniques have shown important drawbacks. Thus, it is reported that thermal techniques require the use of very high temperatures that unfortunately also favor the decomposition rate of the ferrate generated. Likewise, wet or dry oxidation technologies also lead to important drawbacks such as low efficiencies and in some cases the use of hazardous compounds as reagents. In this context, electrochemical synthesis of ferrate using iron electrodes in alkaline solutions has shown better yields, but it has also shown significant problems such as the formation of passivation layers on the electrode surfaces. Figure 12.23 shows the variation of the concentration of ferrate and the current efficiencies during the electrolysis with conductive-diamond anodes of an alkaline solution (14 M NaOH saturated with iron (III) hydroxide). As can be seen, there is a rapid increase in the ferrates concentration during the first stages of the electrolyses. Then, the oxidation rate decreases markedly to a constant value, and the ferrate concentration starts to increase in a slower way. Consequently, the current efficiency of the electrosynthesis decreases during the electrolyses being very low even at the initial stages of the electrolyses. In this context, the maximum concentration of soluble iron species (around 0.2 mM) clearly suggests that the small amounts of available iron can limit markedly the efficiency of the electrosynthesis [54,55].

12.6 SYNTHESIS OF FERRATES

303

Oxidants / mmol FeO42–dm–3

0.07 0.06 0.05 0.04 0.03 0.02 0.01 0 0

1

2

3

4

Q (Ah dm–3)

Current efficiency/%

1 0.8 0.6 0.4 0.2 0

0

1

2

3

4

Q (Ah dm–3) Figure 12.23 Batch production of ferrates with diamond electrodes (C0 :14 M NaOH, j: 1300 A m−2 , T: 30◦ C). (Reprinted from Ref. 54.)

Figure 12.24 shows the effect of the raw matter on the efficiency of the process. Notice that the process depends strongly on the concentration of hydroxyl ions, and it is only efficient for a very high content of this anion. According to literature, the stability of ferrate is greatly influenced by the pH of the system [45,56,57], and ferrate salts seem to be more stable in strongly alkaline conditions. In this context, hydroxide anion concentration was also found to be a key parameter in the electrosynthesis of ferrates with iron electrodes [45]. To increase the availability of oxidizable-iron species, and therefore improve the efficiency of the electrosynthesis with BDD, an iron-powder bed placed near to the anode surface (separated of the anode surface by means of a very thin plastic mesh) was used as raw material. During the electrolysis of aqueous solutions, the oxygen evolution that takes place on the anodic surface, consumes large amounts of hydroxyl anions. This fact leads to changes in the pH in a region very close to the anode surface that can favor the chemical dissolution of the iron particles [58]. This would increase the amount of iron ionic species in the reaction system (coming from the dissolution of iron particles) and, consequently, the efficiencies of the electrosynthesis of ferrate. Part b of Figure 12.24 shows the variation of the ferrate concentration with the electrical charge passed during

304

PRODUCTION OF STRONG OXIDIZING SUBSTANCES WITH BDD ANODES

Oxidants/mmol FeO42–dm–3

0.1

0.08

0.06

0.04

0.02

0 10 5 NaOH(mol dm–3)

0

15

Oxidants/mmol FeO42–dm–3

1.2 1 0.8 0.6 0.4 0.2 0 0

50

100 150 200 250 300 350 400 450 500 550 Q (Ah dm–3)

Figure 12.24 Effect of the raw materials on the production of ferrates with diamond anodes (C0 : 10 M KOH, j: 1000 A m−2 , T: 10◦ C). (Reprinted from Ref. 55.)

the electrolysis of hydroxide solution (10 M KOH) using iron-powder as raw iron. As the figure shows, high ferrates concentrations are obtained (onefold higher than those obtained with Fe(OH)3 ). This also means that efficiencies increase by onefold and the process could become interesting from the commercial viewpoint. Figure 12.25 shows the effect of the operation conditions in electrolysis with the ironpowder bed as raw material for two particular current charge passed. The concentration of obtained ferrates is strongly dependent on the temperature of the electrosynthesis, and the generation process seems to be favored at temperatures around 25◦ C. The observed maximum may be explained in terms of two opposite processes: the solubility of iron (III) (raw material) and the stability of ferrates. Thus, low temperature can lead to lower concentration of iron species available to be oxidized and thus to a more significant masstransfer control. This explains the lower efficiency obtained at this operation condition.

12.7 EFFECT OF THE TYPE OF DIAMOND ON THE EFFICIENCY OF THE PRODUCTION OF OXIDANTS

305

Oxidants / mmol FeO42–dm–3

1.6 1.4 1.2 1 0.8 0.6 0.4 0.2 0 0

1500 500 1000 Current density(A m–2)

2000

Oxidants / mmol FeO42–dm–3

1.6 1.4 1.2 1 0.8 0.6 0.4 0.2 0 0

10

20 30 40 Temperature(°C)

50

60

Figure 12.25 Effect of the operation conditions on the results of the production of ferrates (C0 : 10 M KOH). (Reprinted from Ref. 54.)

Conversely, higher temperatures can disfavor the stability of the oxidant electrogenerated and lead to higher decomposition rate. The effect of current density is as expected according to the results shown for other oxidants produced with conductive diamond. The higher the current density, the better the results obtained. This behavior can be explained in terms [27] of the hydroxyl radical contribution, and indicates that the role of hydroxyl radicals is not only important in the production of peroxo compounds but also on the production of other oxidants. 12.7 EFFECT OF THE TYPE OF DIAMOND ON THE EFFICIENCY OF THE PRODUCTION OF OXIDANTS The rapid development of the diamond technology has focused attention on the search for applications and not on the fundamental aspects of diamond technology. Electrocatalysis of conductive-diamond anodes is still unclear, and the effects of the characteristics of diamond on the process efficiencies have to be further studied in the near future. An example of this is shown in Figure 12.26 for the production of peroxophosphates salts during the discontinuous electrolyses of alkaline solutions containing 1 M K3 PO4 .

306

PRODUCTION OF STRONG OXIDIZING SUBSTANCES WITH BDD ANODES

2.2 Oxidants/mmoles P2O84–

2 1.8 1.6 1.4 1.2 1 0.8 0.6 0.4 0.2 0 0

0.5

1

2

1.5

Q (Ah dm–3) Figure 12.26 Electrolyses of potassium phosphate solution at pH12.5 and current density of 1250 Am−2 . (♦) 1049-4, () 1038-4, (△) 1050-4, (×) 983-5, ( ) 802-3, (◦) 60–61, (+) 908-5, () 858-5, (-) 792-1, ♦ 805-1, (•) 805-3, () 825-6.

Table 12.1 summarizes the main characteristics of the lots of conductive-diamond electrodes used in this work [59,60]. In every case, the pH was controlled at 12.5 ± 0.1, and the current density was fixed at 1250 A m−2 . As it can be observed, the use of conductive diamond with different characteristics leads to a variety of responses in which difference is even more than onefold. However, almost no differences are observed if the experiments are repeated with the same electrodes or with other electrodes of the same lot. This observation confirms that the characteristics of electrodes influence strongly on the electrosynthesis results; hence, they should be taken into account in the study of the applications of conductive-diamond electrolyses. Figure 12.27 shows the results of simple statistical analyses of these results in which the effect of different parameters of the diamond surfaces on the efficiency of production of peroxophosphates is compared. As seen, boron content and thickness of the diamond anode seem to have a significant effect on the production of peroxodiphosphates, ˜ TABLE 12.1 Characteristics of the conductive-diamond electrodes lots used by Canizares et al. [54]. Reference of the BDD 1049-4 1050-4 60-61-G1 1038-4 825-6 908-5 983-5 792-1 805-1/3 858-5 802-3

Conductive-diamond layer Boron Ratio Thickness BDD layer (µm) contents (ppm) sp3 /sp2 100 200 500 1300 1300 2500 1300 100 1300 2500 8000

65 75 93 66 45 68 77 89 105 43 80

1.09 1.14 2.4 1.33 2.33 1.15 2.27 1.03 2.25 1.13 1.05

p-Si substrate Si-Resistivity Roughness, (m cm) Si-Surfinra (µm) 100 100 100 100 100 100 10 10 10 10 10

0.3-0.5 0.3-0.5 0.3-0.5 0.3-0.5