Electrochemical and Electrocatalytic Reactions of Carbon Dioxide Editor-in-Chief: B.R Sullivan Department of Chemistry University of Wyoming Laramie, WY, USA Associate Editors: K.Krist Gas Research Institute Chicago, IL, USA H.E. Guard Office of Naval Research Department of the Navy Arlington, VA, USA 1993 ELSEVIER Amsterdam • London • New York • Tokyo ELSEVIER SCIENCE PUBLISHERS B.V. Sara Burgerhartstraat 25 P.O. Box 211, 1000 AE Amsterdam, The Netherlands Library of Congress Cataloging -in -Publication Data Electrochemical and electrocatalytic reactions of carbon dioxide / edited by B.P. Sullivan, K. Krist, H.E. Guard. P. cm. Includes bibliographical references and index. ISBN 0-444-88316-9 1. Carbon dioxide. 2. Electrochemistry. 3. Catalysis. I. Sullivan B.P. (B. Patrick) II. Krist, K. III. Guard, H.E. QD181.C1E44 1992 546' .6812—dc20 92-34495 CIP ISBN: 0-444-88316-9 © 1993 Elsevier Science Publishers B.V. All rights reserved. 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 or otherwise, without the prior written permission of the publishers, Elsevier Science Publishers B.V., Copyright & Permissions Department, P.O. Box 521, 1000 AM Amsterdam, The Netherlands. Special regulations for readers in the U.S.A. This publication has been registered with the Copyright Clearance Center Inc. (CCC), Salem, Massachusetts. Information can be obtained from the CCC about conditions under which photocopies of parts of this publication may be made in the U.S.A. All other copyright questions, including photocopying outside of the U.S.A., should be referred to the copyright owner, Elsevier Science Publishers B.V., unless otherwise specified. No responsibility is assumed by the publisher for any injury and/or damage to persons or property as a matter of products liability, negligence or otherwise, or from any use or operation of any methods, products, instructions or ideas contained in the material herein. This book is printed on acid-free paper. Printed in The Netherlands. PREFACE The recycling of atmospheric molecules for use as fuels and chemicals is a goal which can be achieved only with a deeper understanding of catalytic processes, particularly electrocatalysis whereby redox transformations can be interfaced with solar or nuclear energy input. Carbon dioxide is a prototypical small molecule in many regards since it is chemically inert. In addition, because of the likely role of carbon dioxide in gobal temperature cycles, it may be that future regulation of output from industrial processes will be desirable. It is the purpose of this volume to present a unified discussion of the carbon dioxide chemistry that is necessary for the understanding and design of electrochemically-driven processes for the reduction of carbon dioxide. The organization of this volume is as follows. First, in chapter 1, a presentation of thermodynamics and kinetics of carbon dioxide reactivity sets the stage for an in-depth discussion of the binding of carbon dioxide to transition metal centers in chapter 2. Chapter 3 presents mechanistic interpretations of perhaps the most important catalytic reaction of C0 , the water-gas shift. In 2 chapter 4 methods for the concentration of C0 , which could prove useful for 2 providing feedstocks for catalytic systems, are presented. Chapter 5 summarizes the state-of-the-art concerning electrochemical mechanisms of C0 reduction at 2 metal centers, focusing principally on the processes that take place in homogeneous solution. Chapters 6 and 7 are devoted to aspects of the reduction of C0 at metal electrodes, both in preparative and mechanistic 2 details. Lastly, chapter 8 covers the more unusual photochemical and photoelectrochemical methods of activating and reducing carbon dioxide. It is hoped that this volume will provide an impetus for further development of electrocatalytic carbon dioxide chemistry. B. Patrick Sullivan Laramie, Wyoming June 1992 LIST OF CONTRIBUTORS W. Bell Solar Energy Research Institute Golden, CO 80401 Ronald L. Cook Eltron Research, Inc. Aurora, IL 60504 Carol Creutz Chemistry Department Brookhaven National Laboratory, Upton NY, 11973 Daniel L. DuBois Solar Energy Research Institute Golden, CO 80401 Peter C. Ford Department of Chemistry University of California Santa Barabara,CA 93111 Karl W. Frese, Jr. Interfacial Sciences, Inc. Santa Clara, CA 95051 F. Richard Keene James Cook University of North Queensland Department of Chemistry and Biochemistry Townsville, Queensland 4811, Australia Nathan S. Lewis Department of Chemistry and Chemical Engineering, California Institute of Technology, Pasadena, CA A. Miedener Solar Energy Research Institute Golden, CO 80401 Anthony F. Sammells Eltron Research, Inc. Aurora, IL 60504 Gary A. Shreve Department of Chemistry and Chemical Engineering, California Institute of Technology, Pasadena, CA J. C. Smart Solar Energy Research Institute Golden, CO 80401 B. Patrick Sullivan Department of Chemistry, University of Wyoming, Laramie, WY 82071-3838 1 Chapter 1 THERMODYNAMIC, KINETIC, AND PRODUCT CONSIDERATIONS IN CARBON DIOXIDE REACTIVITY F. Richard Keene Department of Chemistry and Biochemistry, James Cook University of North Queensland, Townsville, Australia 4811 There has been a recent upsurge in interest in the reactivity of carbon dioxide for two primary reasons. Firstly, carbon dioxide is the ultimate by-product of all processes involving oxidation of carbon compounds and its increasing presence in the atmosphere since the beginning of the Industrial Revolution has given rise to widespread concern about possible consequences (the so- called "Greenhouse Effect"). Secondly, in view of the vastness of its supply, carbon dioxide represents a possible potential source for Q feedstocks for the manufacture of chemicals and fuels, alternative to the current predominant use of petroleum-derived sources. Carbon reserves in the form of atmospheric carbon dioxide, carbon dioxide in the hydrosphere and carbonates in the terrestrial environment substantially exceed those of the fossil fuels such as coal and petroleum (1,2). Of course, the conversion of carbon dioxide to fuels and chemicals is carried out in the biosphere on an enormous scale by photosynthesis. The raising of consciousness on the industrial/ecological issues of limitations to fossil fuel reserves and of the consequences of their extensive use has heightened interest in the concept of recycling carbon resources, and therefore in the activation of carbon dioxide and "artificial photosynthesis" (3). Many of the aspects of the thermodynamics of carbon dioxide reactivity, as well as considerations of carbonate chemistry, were extensively reviewed in 1983 by Palmer and van Eldik (4). The present Chapter seeks to present a brief overview of some of the properties and reactivity of carbon dioxide, and to highlight potential means of promotion of its reactivity: many of the issues raised will be elaborated upon in subsequent chapters in this volume. 1. Physical Properties of Carbon Dioxide 1.1 Structure and Bonding. Carbon dioxide is a linear molecule for which the following canonical structures can be drawn: + . . .. + :0 C = 0: o : 0 = C = 0: <-» :0=C 0 = Despite the linear symmetry and overall nonpolar nature of the molecule, some chemical reactivity might be anticipated associated either with the presence of the 7t-electron density of the double bonds and the lone pairs of electrons on the oxygen atoms, or with the electrophilic carbon atom. The 2 qualitative MO energy level diagram is given in Figure 1, and the estimation of the level of the lowest unoccupied antibonding orbital {ca. 3.8 eV (1)} indicates a high electron affinity associated with the central carbon atom. C orbltals O orbitals Figure 1. MO energy level diagram for the linear triatomic molecule CO2 On the other hand, the first ionization potential is high {13.7 eV (1)}, so that the electrophilicity of the central carbon atom might be anticipated as the site of predominant reactivity. 1.2 Stability, Carbon dioxide, CO2, is the ultimate product of oxidation of carbon and its compounds and is extremely thermodynamically stable {AG° = -394 kJ mol1, cf. -137 kJ mol1 for carbon monoxide, CO (5,6)}. The bond strength of the C-0 bond in carbon monoxide is the largest known {D = 1076 kJ mol1 (7)}. In carbon dioxide, that bond strength is measured at D = 532 kJ mol-1 (7). 1.3 Carbon Dioxide Reduction. The redox potentials for the reduction of carbon dioxide have been determined: 3 C0 <=* C0- E° =-1.90V (8),-1.85V {vs NHE(9)} 2 2 C0- <=> C02" E° = -1.2V{vs NHE(IO)} 2 2 For the first reduction, there is a change in geometry from the linear CO2 to a bent CO2"* (11): this structural change gives rise to a very slow self-exchange rate for the CO2/CO2" couple (12) and to a significant overpotential in the reduction of CO2 (13). As noted previously by Schwarz et al (12), the nature of CO2"* may be significant in the consideration of aspects of the activation of CO2. 2. Organic Reactions of Carbon Dioxide Carbon dioxide has a reasonably extensive organic chemistry, which has been reviewed previously (14-16). Much of this chemistry requires catalysis in some form, and reactions are generally represented within one of the following categories - • reductions of CO2 by H2: Via the water-gas-shift reaction (WGS - equation (1)) C0 + H ?=± CO + H 0 [1] 2 2 2 and technologies such as the Fischer-Tropsch processes (17), a variety of products may be produced including methanol, alkanes, alkenes, alcohols, ethers, esters, aldehydes, ketones, etc. ♦ carboxylation of active hydrogen compounds: e.g. carbon dioxide reacts with alkynes to produce pyrones, with 1,3-butadienes and allenes to produce esters and lactones, with methylenecyclopropanes to produce lactones, and with epoxides to produce organic carbonates. 4 chemistry related to the Kolbe-Schmitt reaction (i.e. the reaction of CO2 with phenolates): e.g. alkali metal phenolates react with CO2 to produce 2- and 4-carboxylated phenols. OK OK OH CO, COOK the reaction of CO2 with amines (including ammonia): e.g. CO2 undergoes reaction with primary and secondary amines to yield carbamates (which may react further in the presence of appropriate substrates), and with ammonia under forcing conditions to produce urea. R NH + C0 ^ R2NCOOH 2 2 2NH + C0 -> H NC(0)NH + H 0 3 2 2 2 2 Within this diversity, there are only four major industrial chemical processes which use carbon dioxide as a carbon feedstock. (1) H2 (WGS) (2) Fischer-Tropsch synthesis Alkanes OH Alkenes Alcohols Aldehydes Ketones Esters C02Na Ethers The commercial use of CO2 is quite extensive, particularly in refrigeration and as a cryogen, in fire extinguishers and in beverages. The reduction reactions involving molecular hydrogen are instructive in terms of the thermodynamics of carbon dioxide reactivity. The free energies on reduction of C02() by molecular g hydrogen, and in a limited number of organic reactions, are calculated and shown in the Table below (5,6). Clearly, the thermodynamic stabilities of the CO2 and H2O molecules are dominant in the consideration of the reactions of CO2 with molecular hydrogen. In cases where water acts as an "oxygen sink" the chances of a thermodynamically favorable reaction are higher. However as has been pointed out previously (16,18), molecular hydrogen is produced industrially by the involvement of the water-gas-shift reaction {equation (1)}, whence CO2 itself is the ultimate "oxygen sink". This irony illustrates the necessity for alternative reduction strategies for carbon dioxide reactivity, and carbon dioxide reduction in particular, among which the reaction of CO2 with metal-hydride species must be of fundamental interest. REACTION AG°(kJmol-0 1 C0 + H -> CO + H0i) +19.9 2 2 2 ( C0 + H -> HCOOH +48.4 2 2 (1) C0 + 2H -> HCHO(g) + H0 +47.2 2 2 2 C0 + 3H -> CHOH + H0 -9.1 2 2 3 (1) 2 C0 + 4H -> CH + 2H0 -130.8 2 2 4 2 2C0 + H -> (COOH) ) +90.9 2 2 2(s 2C0 + 6H -> CHOCH() + 3H0 -36.8 2 2 3 3g 2 C0 + CH(i) -> CHCOOH +18.7 2 6 6 6 5 (s) C0 + CH -» CHCOOH +53.2 2 4 3 (1) C0 + CH + H -> CHCHO + H0 +74.4 2 4 2 3 (g) 2 C0 + H + CHOH -* CHCOOH(i) + H0 -68.5 2 2 3 (1) 3 2 C0 + 3H + CHOH -> CHCHOHi) + 2H0 -88.5 1 2 2 3 (1) 3 2 ( 2 These values can be compared with analogous reactions for carbon monoxide: e.g. REACTION AG^kJmol1) 1 CO + 2H -> CHOHi) -29.0 2 3 ( CO + H0 -> HCOOH +28.5 2 (1) CO + CHOHi) -^ CHCOOHi) -88.4 1 3 ( 3 ( 3. Biological Reactions of Carbon Dioxide Carbon dioxide is nature's primary source of carbon, particularly through photosynthetic fixation to form carbohydrates. In terms of other chemical reactions of C0 in biological systems, 2 many may be broadly classified as carboxylation reactions involving the electrophilic addition of CO2 to a substrate anion, as for example in the enzymic conversion of phosphoenolpyruvate to oxalacetate: H2C =^ RX "CO, coo- RXPO02' "OOC COO" (where RX is a phosphoryl acceptor such as water, a nucleoside diphosphate, or inorganic phosphate). In addition, the hydration/dehydration of C0 is essential in respiration and is catalyzed 2 by the enzyme carbonic anhydrase. C02 H20 H2CO3 HCO3- H+ A discussion of the variety of these biological reactions is not warranted here, but they have been reviewed elsewhere (19). 6 4. Equilibria and Kinetics of Aqueous Solutions Containing Carbon Dioxide Many of the experimental studies addressing aspects of the solubility of carbon dioxide in a variety of solvents, and of the nature and kinetics of the resultant equilibria, have been extensively reviewed (4) and only a summary of some of those details will be discussed here. On dissolution of gaseous carbon dioxide in water, a rapid C02() <^ CC>2(aq) equilibrium g occurs, whereupon a slow equilibrium (equation (2)) is established between loosely hydrated CCfyaq) and "carbonic acid", H2CO3. ki C0 (aq) + H 0 +± H2CO3 (K ) [2] 2 2 H k-1 The equilibrium constant, KH, for this reaction can be calculated as 2.6 x 103 at 25 °C: this value shows only minor temperature variation. Average values of ki and k_i (25°C; zero ionic strength) are 6.2 x 10-2 s_1 and 23.7 s_1, respectively (4). The ionization constant for the dissociation of "carbonic acid" H2CO3 <=* H+ + HCCV (K ) [3] 0 has been determined (20,21) to be 1.7 x 10~4 (25°C), and the reactions are very rapid (the rate of protonation of HCO3- is virtually diffusion controlled (22)). Accordingly, although the first ionization to produce the bicarbonate ion, HCO3-, is better expressed in the form C0 (aq) + H 0 <=* H+ + HCO3" (K0 [4] 2 2 since <1% of the dissolved CO2 is present as H2CO3, H2CO3 is an intermediate in reaction (3). Ki is often referred to as the "apparent" acid dissociation constant of H2CO3, and should have a value (KOKH) of ca. 4 x 10-7 under the above conditions. At higher pH values, CCtyaq) may react directly with OH' C0 q) + OH" T± HCO3- (Ki) [5] 2(a where the forward and back reactions determined at 7.7 x 103 M_1s_1 (23,24) and 2.3 x 10-4 s-1 (25) at 25°C (respectively) lead to K « 3 x 107 M1. r The second dissociation reaction, viz. HCO3- ^ H+ + CO32- (K ) [6] 2 appears normal in every respect and K2 = 4.7 x 1011 at 25°C (26). A further equilibrium, called "carbonate catalysis" may also be present in solutions containing CO32 : C0 q) + CO32- + H 0 *± 2HC0 - [7] 2(a 2 3 however, unless a buffer containing carbonate ion were being used, rate studies indicate the contribution of such a path to be negligible by comparison with the previous equilibria (4,27).