摘要:

PARADAY DISCUSSIONS OF THE CHEMICAL SOCIETY NO. 56 1974 In termed a tes in Electro- chemical Reactions THE FARADAY DIVISION CHEMICAL SOCIETY LONDONISBN: 0 85186 808 2 Library of Congress Catalog Card No. 72-92915A GENERAL DISCUSSION ON Intermediates in Electrochemical Reactions 18 t h-20 th Sep tern ber, 1973 A GENERAL DrscussioN on Intermediates in Electrochemical Reactions was held at the University of Oxford on the 18th, 19th and 20th September 1973. The Presi- dent, Prof. Sir George Porter, F.R.S., was in the Chair at the first session and 169 members and others were present. Among the overseas visitors were: Dr. N. M. Alpatova, U.S.S.R. Prof. M. M. Baizer, U S A . Dr. A. J. Bard, France Prof. E. Barendrecht, The Netherlands Mr. G. Bech-Nielsen, Denmark Dr.G. Bottura, Italy Dr. D. J. Brown, W. Germany Prof. S . Bruckenstein, U.S.A. Mr. H. Bruder, W. Germany Dr. L. D. Burke, Eire Prof. G. Cauqius, Fruncc Prof. B. E. Conway, Canada Dr. A. Damjanovic, U.S.A. Prof. X. De Hemptinne, Belgium Dr. W. J. De Klein, The Netherlands Prof. A. Despic, Yugoslavia Dr. J. Dunnett, Switzerland Prof. I. Epelboin, France Mr. K.-D. Franz, W. Germany Dr. C . L. Gardner, Canada Prof. H. Gerischer, Germany Prof. E. Gileadi, Israel Dr. S. Gottesfeld, Israel Dr. A. D. Grishina, U.S.S.R. Dr. M. M. P. Janssen, The Netherlilnds Dr. J. 1. Japaridze, U.S.S.R. Mr. B. S. Jensen, Denmark Dr. B. Kastening, W. Germany Mr. M. Keddam, France Dr. R. Knoedler, W. Germany Dr. D. M. Kolb, Germany Dr. H. A. Kozlowska, Cunadu Prof. T. Kuwana, U.S.A. Dr.J. C. Lestrade, France Prof. P. Martinet, France Dr. J. D. E. McIntyre, U.S.A. Dr. W. Mehl, Switzerland Dr. E. Nicolas, Belgium Prof. J. Padova, Israel Prof. G. Pilloni, Ztaly Dr. Yu. V. Pleskov, U.S.S.R. Dr. L. Pospisil, Czechoslovakiu Dr. V. Pravdic, Yugoslavia Dr. J. C . Reeve, Denmark Prof. R. E. Rieke, U.S.A. Dr. J. Riha, Czechoslovakiu Dr. S. Roffia, Italy Dr. D. Serve, France Dr. D. J. Schiffrin, Aryentinu Prof. J. W. Schultze, Germany Dr. R. E. Sioda, Poland Dr. F. R. Smith, Canadu Mr. A. Spaans, The Netherhiids Dr. R. Tomat, Italy Prof. J. W. Tomlinson, New Zeuland Mr. J. Tenygl, Czechoslovakia Dr. S . Vdkher, Ituly Prof. N. Vandenborgh, U.S.A. Dr. D. Van der Meer, The Netherhiids Mr. D. Vattis, Greece Prof. K. Vetter, Germany Mr. E. Vieil, France Prof.A. A. Vlcek, Czechoslovukia Dr. J. Volke, Czechoslovakia Mr. M. Vukovic, Yicgoslavia Dr. J. A. Wargon, Argentinu Dr. J. Weber, Czechoslovakicr0 The Chemical Society and Contributors 1974 Printed in Great Britain at the University Press, AberdeenCONTENTS page 7 Introduction by H. Gerischer 16 Spectroelectrochemical Methods in the Study of Short-lived Intermediates by R. F. Broman, W. R. Heinenian and T. Kuwana 28 41 52 62 75 96 108 122 138 152 171 180 199 Photochemical Generation of Semiquinone Intermediates at a Rotating Semi- transparent Disc Electrode by W. J. Albery, M. D. Archer, N. J. Field and A. D. Turner Charge Transfer Reactions involving Intermediates formed by Homogeneous Capture of Laser-produced Photoelectrons by G. C. Barker, D. McKeown, M.J. Williams, G. Bottura and V. Concialini Investigation of Intermediates by Electron Photoemission from Metal itito Electrolyte Solution by Yu. V. Pleskov, Z . A. Rotenberg, V. V. Eletsky and V. I. Lakomov GENERAL DIscussroN.-Dr. R. Parsons, Prof. A. A. VIEek, Prof. Sir George Porter, Prof. T. Kuwana, Dr. A. Bewick, Dr. B. Kastening, Dr. W. J. Albery, Dr. B. R. Eggins, Dr. R. M. Reeves, Dr. G. C. Barker, Dr. Frank R. Smith, Dr. Yu. V. Pleskov, Dr. Z. A. Rotenberg, Dr. V. V. Eletsky, Dr. V. I. Lakomov, Prof. J. W. Schultze, Prof. I. E. Epelboin, Dr. D. J. Schiffrin Application of the Photo-electrochemical Eflect to the Stud)) of the Electro- chemical Properties of Radicals: CO; and CH;, by D. J. Schiffrin Studies of Electrochemically Generated Reaction Intermediates using Modu- lated Specular Reflectance Spectroscopy by A.W. B. Aylmer-Kelly, A. Bewick, P. R. Cantrill and A. M. Twxford GENERAL DIscussIox-Dr. W. J. Albery, Dr. B. R. Eggins, Dr. D. Schiffrin, Dr. A. Bewick, Dr. 0. R. Brown, Dr. R. M. Reeves, Dr. G. C . Barker, Dr. R. Parsons, Dr. Frank R. Smith Prof. B. E. Conway Optical Spectroscopy of Adsorbed Intermediates in Electrochemical Reactions by J. D. E. McIntyre and W. F. Peck, Jr. D i e r en t ial Reflectance Spectroscopy of Mono layer an d Sii bm on o layer Met a1 Deposits on Semiconductor and Metal Electrodes by D. M. Kolb GENERAL DwxssIoN.-Dr. I. Bergmail, Prof. B. E. Conway, Dr. A. Bewick, Dr. T. Kuwana, Dr. J. D. E. McIntyre, Dr. C . L. Gardner, Dr. E. J. Casey, Miss M. A. Barrett, Dr. H. Angerstein-Kozlowska, Dr.D. M. Kolb, Dr. M. Keddam, Dr. A. J. McQuillan, Dr. J. S. Clarke, Dr. A. T. Kuhn, Prof. W. J. Orville-Thomas Intermediates in Electro crys tallisa t ion by W. Davison, J. A. Harrison and J. Thompson Nucleation by P. Bindra, M. Fleischmann, J. W. Oldfield and D. Singleton GENERAL DIscussIoN.-Dr. R. M. Reeves, Dr. J. A. Harrison, Prof. A. R. DespiC, Prof. A. A. VlEek, Dr. M. VukoviC, Dr. V. PravdiC, Prof. I. Epelboin, Dr. M. Keddam, Prof. M. Fleischmann, Prof. M. R. Thirsk, Dr. D. J. Schiffrin, Prof. H. Gerischer 56 210 228 235 244 264 276 285 293 308 317 330 341 353 367 379 CONTENTS Kinetic and Optical Relaxation Studies of Adsorbed Intermediates in Some Electrochemical Reactions by B. E. Conway, H. Angerstein-Kozlowska, F. C. Ho, J. Klinger, B.MacDougaIl and S. Gottesfeld A Combined Adsorption Isotherm for Intermediates formed in Electrode Reactions by E. Gileadi GENERAL D~scuss~o~.-~rof. B. E. Conway, Dr. A. Capon, Miss M. A. Barrett, Prof. J. W. Schultze, Prof. K. J. Vetter, Dr. 0. R. Brown, Dr. M. Keddam, Dr. B. Kastening, Dr. R. Parsons, Prof. E. Gileadi, Dr. H. P. Dhar The A . C. Impedance of Complex Electrocliemical Reactions by R. D. Armstrong, R. E. Firman and H. R. Thirsk Faradaic Impedances and Intermediates in Electrochemical Reactions by I. Epelboin, M. Keddam and J. C. Lestrade Detection of Electrochemically Generated Intermediates by Variable Sweep Rate Cyclic Voltammetry by B. R. Eggins Electrochemical Mass Spectrometry. I. Preliminary Studies of Propane Oxidation on Platinum by S. Bruckenstein and J.Comeau GENERAL D~scussro~.--Dr. L. PospiSil, Dr. M. Keddam, Prof. M. Fleisch- mann, Prof. I. Epelboin, Dr. J. C. Lestrade, Dr. B. R. Eggins, Dr. R. M. Reeves, Dr. Kuwana, Prof. B. E. Conway, Dr. J. Volke, Dr. 0. Manousgk, Prof. A. A. VlEek, Dr. A. Rusina, Dr. J. Heitbauni, Dr. D. J. Brown Detection and Identgcation of Intermediate and Final Products of Electro- chemical Reactions by Means of the Rotating Ring- Disc Electrode Method by L. N. Nekrasov Alternating Current and Ring-Disc Electrodes by W. J. Albery, A. H. Davis and A. J. Mason Simultaneous Electrochemical and E.s.r. Studies on the Reduction of A lipha t ic Nitr ocompo unds by J. K. Dohrmann, F. Gallusser and H. Wittchen E.s.r. Investigations of Short-lived Intermediates in Electrochemical Reac- tions by B. Kastening, B. GostiSa-MihelCiC and J. DiviSek Reductive Coupling and Isomerization of Electrogenerated Radical Ions of cis- and trans-Isomers by A. J. Bard, V. J. Puglisi, J. V. Kenkel and A. Lomax GENERAL Drscuss~o~.-~rof. A. R. DespiC, Prof. I. Epelboin, Prof. H. Gerischer, Dr. M. Keddam, Prof. S . Bruckenstein, Dr. W. J. Albery, Prof. A. J. Bard, Prof. J. K. Dohrmann, Dr. E. J. Casey, Dr. C. L. Gardner, Dr. B. Kastening, Prof. M. Fleischmann, Dr. R. E. W. Jansson, Prof. A. J. Bard, Prof. R. D. Rieke Concluding Remarks by J. E. B. Randles 383 AUTHOR INDEX

ISSN:0301-7249    

DOI:10.1039/DC9735600001

出版商:RSC    

年代:1973

数据来源: RSC

摘要:

GENERAL DlSCUSSlONS OF THE FARADAY SOCIETY Date 1907 lW7 1910 1911 1912 1913 1913 1913 1914 1914 1915 1916 AYlb 1917 1917 1917 1918 1918 1918 1918 1919 1919 192U AYW AY2U 192u AYU 1921 192l 1921 A922 1922 1923 1923 1923 1923 19L3 I924 1924 1924 1924 1924 1925 1925 1926 1926 1927 1927 1927 Subject 08mOtic prcaure Hydrates in Solution Tho Constitution of Water Hioh Tomporature Work Magnew Properties of Alloy Colloid8 and thoir VLcority The Corrorion of iron and Stml The P-vity of Motah OpUcal Rotary P o w Tbs Hardemn ' gof Metab The Tramformation of Pure Iron Method8 and Appliances for tbo Attainment of High Temperaturea in 8 Labomtory Rsfractory Matedab Training and Work of tho cbamical Engineer olmotic Pruuw Pyromctcm and PyronuW Tho Setting of camsllu and Plrrtan E l a c t r i c a l F m Ceordinatioa of Scientific Publication Th Occlusion of Gaaa by Metab Tba Prosont Position of the Thaory of Ionization Tho Expmination of Matmiah by X-Ray8 Tbu Miaoscopa : It, DeiQp, Comtruction and Applications Basic Slags: Tbeir Production and Utilization in Agriculture Physica and C b d a t f y of C O W Ek#r&@tion and Ebctroplating crpiuanv Tb6 Failw of Metab under Internal and Prolonged Sttess Phydco-Chemid Probbmr Relating to the Soil catalysib with dbmnoe to Newer Thaorica of Chemical Ador, Some Propmtia of Powdm with special reference to Gradha by Hutriation The Generation and Utilization of Cold Alloys Resistant to Conrosion The Physical -try of tha Photographic ROCUU The Uccuonic Theory of Valcncy Electrode Rcactiolu and Equilibria Atmospheric Corrosion.First Report Investigation on Oppau Ammonium Sulphatc-Nitrat+ Fluxes and Slags in Metal Melting and Working Physical and Physico-cherm 'cal problems dating to Textile F i b The Physical Chemistry of lgnwus Rock Formation The Physical Chtmietry of Steel-Making Processes Photochemical Reactions in Liquids and Gascs Exprosive Reactions in Gaseous Media Physical Phenomena at interfacc~, with s@al reference to MOkdW A m ~ @ ~ i c Corrosion. Second Report Tbe Thaory of Strong ElactrolW Cohesion and Rdatbd ProblemS Base Exchange in soils -tation volvnr m. 3 3 6 7 8 9 9 9 10 10 11 12 12 13 13 13 14 14 14 14 15 15 16 16 16 16 17 17 17 17 18 18 19 19 19 19 19 20 20 20 20 20 21 21 22 22 23 23 2AGENERAL DISCUSSIONS OF THE PARADAY SOCIETY Date 1928 1929 1929 1929 1930 1930 1931 1932 1932 1933 1933 1934 1934 1935 1935 1936 1936 L 937 1937 1938 1938 1939 1939 1940 1941 1941 1942 1943 1944 1945 1945 1946 1946 1947 1947 1947 1 947 1948 1948 1 949 1949 1 949 1950 1950 1950 1950 1951 1951 1952 1952 1952 1953 1953 1954 1954 Subject Homogeneous Catalysis Crystal Structure and Chemical Constitution Atmospheric Corrosion of Metals.Third Report Molecular Spectra and Molecular Structure Optical Rotatory Power Colloid Science Applied to Biology Photochemical Procases Tho Adsorption of Gaws by Solids The Colloid Aspect8 of Textile Materials Liquid Cry8tah and Anisotropic Melts Free Radicals Dipob Moments Colloidal Electrolytes Tho Structure of Metallic Coatings, Films and Surfacea Tho Phenomena of PolymeriZation and Condensation Dirpersa Systems in Gases : Dust, Smoke and Fog Structuro and Molccu€ar Forcas in (a) Pure Liquids, and (b) Solutions The Properties and Functions of Membranes, Natural and Artificial Reaction Kioeticr Chomicd Reactions Involving Solida Hydrocarbon Chembtry The Elactrical Double Layer (owing to the outbreak of war tho meeting Tho Hydrogen Bond Tho Oil-Water Intstfaca ?'ho Mechanirm and Chomical Kinerics of Organic W o r n in Liquid The Structure and Reactions of Rubbar M o k of Drug Action Molecutas Weight and Molecular Weight Distribution in High Polymclr.(Joint Meeting with the Plastics Group, Society of Chemical Industry) The Application of Ma-red Spectra to Chemical Problsmrr Oxidation Dielcctrica Swelling and Shrinking Elactroda Prowsscs The Labile Molecule Surface Chemistry. (Jointly with the Socikt6 de Chimie Physique at was abandonad, but the papers w m printed in tho Transoctlonr) sy8t8llM Bordeaux.) Published by Butterworths Scientific Publications, Ltd.Voliinie 24 25 25 25 26 26 27 28 29 29 30 30 31 31 32 32 33 33 34 34 35 35 35 36 37 37 38 39 40 41 42 42 A 42 B Disc. 1 2 Colloidal Electrolytes and Solutions Traas. 43 The Interaction of Water and Porous Materials Disc. 3 4 Lipo-Proteins 6 Hcterogmeous Catalysis 8 PhysicoGhemical Properties and Behaviour of Nuclear Acids Trans. 46 Spectroscopy and Molecular Structure and Optical Methods of In- vestigating Cell Structure Disc. 9 Electrical Double Layer Trans. 47 Hydrocarbons Disc. 10 The Physical Chemistry of Process Metallurgy crystal Growth 5 Chromatographic Analysia 7 The Size and Shape Factor in Colloidal Systems Radiation Chemistry 12 11 The Physical Chemistry of Proteins 13 The Reactivity of Free Radicals 14 The Equilibrium Properties of Solutions of Non-Elcctrolytca 15 The Physical Chemistry of Dyeing and Tanning 16 The Study of Fast Reactions 17 Coagulation and Flocculation 18GBN5RAL DISCUSSIONS OF THE FARADAY SOCIETY Date 1955 1955 1956 1956 1957 1958 1957 1958 1959 1959 1960 1960 1961 1961 1962 1962 1963 1963 1964 1964 1965 1965 1966 1966 1967 1967 1968 1968 1969 1969 1970 1970 1971 1971 1972 1972 1973 1973 Sdject Volunic Microwave and Radio-Frequency Spectroscopy Physical Chemistry of Enzymes Membrane Phenomena Physical Chemistry of Processes at High Pressures Molccdar Mechanism of Rate Processes in Solids Interactions in Ionic Solutions Configurations and Interactions of Macromolecules and Liquid Crystals Ions of the Transition Elements Energy Transfer with special reference to Biological Systems Crystal Imperfections and the Chemical Reactivity of Solids Oxidation-Reduction Reactions in Ionizing Solvents The Physical Chemistry of Aerosols Radiation Effects in Inorganic Solids The Structure and Properties of Ionic Melts Inelastic Collisions of Atoms and Simple Molecules High Resolution Nuclear Magnetic Resonance The Structure of Electronically-Excited Spec& in tho Gas-Phase Fundamental Processes in Radiation Chemistry Chemical Reactions in the Atmosphere Dislocations in Solids The Kinetics of Proton Transfer Processes Intermolecular Form The Role of the Adsorbed State in Heterogeneous Catalysis Colloid Stability in Aqueous and Non-Aqueous Media The Structure and Properties of Liquids Molecular Dynamics of the Chemical Reactions of Gases Electrode Reactions of Organic Compounds Homogeneous Catalysis with Special Reference to Hydrogenation and Bonding in Metallo-Organic Compounds Motions in Molecular Crystals Polymer salutions The Vitreous State Electrical Conduction in Organic Solids Surface Chemistry of Oxides Reactions of Small Molecules in Excited States The Photoelectron Spectroscopy of Molecules Molecular Beam Scattering Intermediates in Electrochemical Reactions Oxidat ion For current availability of Discwwn volumes, see back cover.15 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56

ISSN:0301-7249    

DOI:10.1039/DC973560X003

出版商:RSC    

年代:1973

数据来源: RSC

摘要:

GBN5RAL DISCUSSIONS OF THE FARADAY SOCIETY Date 1955 1955 1956 1956 1957 1958 1957 1958 1959 1959 1960 1960 1961 1961 1962 1962 1963 1963 1964 1964 1965 1965 1966 1966 1967 1967 1968 1968 1969 1969 1970 1970 1971 1971 1972 1972 1973 1973 Sdject Volunic Microwave and Radio-Frequency Spectroscopy Physical Chemistry of Enzymes Membrane Phenomena Physical Chemistry of Processes at High Pressures Molccdar Mechanism of Rate Processes in Solids Interactions in Ionic Solutions Configurations and Interactions of Macromolecules and Liquid Crystals Ions of the Transition Elements Energy Transfer with special reference to Biological Systems Crystal Imperfections and the Chemical Reactivity of Solids Oxidation-Reduction Reactions in Ionizing Solvents The Physical Chemistry of Aerosols Radiation Effects in Inorganic Solids The Structure and Properties of Ionic Melts Inelastic Collisions of Atoms and Simple Molecules High Resolution Nuclear Magnetic Resonance The Structure of Electronically-Excited Spec& in tho Gas-Phase Fundamental Processes in Radiation Chemistry Chemical Reactions in the Atmosphere Dislocations in Solids The Kinetics of Proton Transfer Processes Intermolecular Form The Role of the Adsorbed State in Heterogeneous Catalysis Colloid Stability in Aqueous and Non-Aqueous Media The Structure and Properties of Liquids Molecular Dynamics of the Chemical Reactions of Gases Electrode Reactions of Organic Compounds Homogeneous Catalysis with Special Reference to Hydrogenation and Bonding in Metallo-Organic Compounds Motions in Molecular Crystals Polymer salutions The Vitreous State Electrical Conduction in Organic Solids Surface Chemistry of Oxides Reactions of Small Molecules in Excited States The Photoelectron Spectroscopy of Molecules Molecular Beam Scattering Intermediates in Electrochemical Reactions Oxidat ion For current availability of Discwwn volumes, see back cover.15 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56GBN5RAL DISCUSSIONS OF THE FARADAY SOCIETY Date 1955 1955 1956 1956 1957 1958 1957 1958 1959 1959 1960 1960 1961 1961 1962 1962 1963 1963 1964 1964 1965 1965 1966 1966 1967 1967 1968 1968 1969 1969 1970 1970 1971 1971 1972 1972 1973 1973 Sdject Volunic Microwave and Radio-Frequency Spectroscopy Physical Chemistry of Enzymes Membrane Phenomena Physical Chemistry of Processes at High Pressures Molccdar Mechanism of Rate Processes in Solids Interactions in Ionic Solutions Configurations and Interactions of Macromolecules and Liquid Crystals Ions of the Transition Elements Energy Transfer with special reference to Biological Systems Crystal Imperfections and the Chemical Reactivity of Solids Oxidation-Reduction Reactions in Ionizing Solvents The Physical Chemistry of Aerosols Radiation Effects in Inorganic Solids The Structure and Properties of Ionic Melts Inelastic Collisions of Atoms and Simple Molecules High Resolution Nuclear Magnetic Resonance The Structure of Electronically-Excited Spec& in tho Gas-Phase Fundamental Processes in Radiation Chemistry Chemical Reactions in the Atmosphere Dislocations in Solids The Kinetics of Proton Transfer Processes Intermolecular Form The Role of the Adsorbed State in Heterogeneous Catalysis Colloid Stability in Aqueous and Non-Aqueous Media The Structure and Properties of Liquids Molecular Dynamics of the Chemical Reactions of Gases Electrode Reactions of Organic Compounds Homogeneous Catalysis with Special Reference to Hydrogenation and Bonding in Metallo-Organic Compounds Motions in Molecular Crystals Polymer salutions The Vitreous State Electrical Conduction in Organic Solids Surface Chemistry of Oxides Reactions of Small Molecules in Excited States The Photoelectron Spectroscopy of Molecules Molecular Beam Scattering Intermediates in Electrochemical Reactions Oxidat ion For current availability of Discwwn volumes, see back cover.15 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56

ISSN:0301-7249    

DOI:10.1039/DC97356BX006

出版商:RSC    

年代:1973

数据来源: RSC

摘要:

Intermediates in Electrochemical Reactions ? BY HEINZ GERISCHER Fritz-Haber-Tnstitut der Max-Planck-Gesellschaft, Berlin-Dahlem Received 20th October, 1973 The search for intermediates in chemical reactions is like a detective story. The kineticist knows the actors, present at the beginning, and he sees the victims. He has to search for witnesses to reconstruct the course of the events. His witnesses are the intermediates. Without such witnesses, his reasoning depends on circumstantial evidence and may often be quite erroneous. It is obvious, therefore, how important such intermediates are for the understanding of electrode reactions and how necessary it is to find the most direct evidence for their existence. However, to introduce such a subject is like trying to achieve an essential synthesis of all detective stories within half an hour.This introduction will therefore be very fragmentary and, necessarily, will reflect a great deal of personal choice. In the history of electrode reactions, one particular intermediate is of extraordinary interest. It is the H-atom, adsorbed-or better expressed chemisorbed-on the surface of a metal. This intermediate plays a key role in the kinetics of the hydrogen evolution and also in the mechanism of electrochemical hydrogenation. It was first introduced by J. Tafel in 1900 as a probable intermediate in cathodic hydrogen evolution, and its properties are still discussed as the papers given by Barker and by Pleskov show in this Discussion. Tafel’s reasoning for the existence of this inter- mediate was a typical example of his use of circumstantial evidence.His principal reason was chemical intuition. He gave, however, a derivation of the (current, voltage) curve based on the kinetics of consecutive reactions, assuming that the recombination of H-atoms would be rate determining.2 This was the kinetic founda- tion of the so-called “ Tafel ” plot, U = a+b log i, widely used in electrochemistry. Although he realised that the theoretical derivation gave a b-value of 30 mV/log i, in contrast to his experimental findings of 120 mV/log i for mercury and lead, he trusted his chemical intuition and he was right in principle, as we know today. He suffered only the misfortune that these electrode materials were not suitable for his mechanistic assumptions. Tafel’s approach served as a model for many other electrode reactions where chemical intuition tells us that the process should occur in consecutive steps.This is especially reasonable for multielectron redox reactions, since simultaneous transfer of several electrons is extremely unlikely. Redox reactions with organic molecules are, therefore, a wide field for the search for intermediates because unstable radicals are the primary product of a one-electron transfer between an electrode and a stable organic molecule. The same is true for all other reactions, where the products differ widely in chemical composition from the initial reactants. However, chemical intuition can be a questionable guide. Not only does it change with increasing experience and theoretical understanding, but also it will quite often present alternatives for the reaction paths.Therefore, it remains the task -1- Introductory Lecture. 7 If he is lucky he can find some.8 INTERMEDIATES I N ELECTROCHEMICAL REACTIONS of the kinetical investigation to find out which path is dominant and which inter- mediates are really formed in a particular reaction. In the discussion which follows I want first to classify the most common types of intermediates in electrode reactions. Secondly, I shall give a characterization of the methods which are available and have been used in finding intermediates with the main emphasis on present developments.~’ CLASSIFICATION OF INTERMEDIATES IN ELECTRODE REACTIONS In my opinion, it seems useful to distinguish between 3 types of intermediates, classified according to their interaction with the surface.Three classes defined in this way are summarized in table 1 . This classification is certainly not unambiguous. TABLE 1 .-CLASSIFICATIONS OF INTERMEDIATES IN ELECTRODE REACTIONS class interaction with electrode predominant location reactions involved 1 no or weak 2 strong solution, to some generated in horno- extent on surface geneous or hetero- geneous reactions interface, to some generated primarily extent in solution in heterogeneous reactions 3 very strong chemical surface only generated in surface bond as part of the solid reactions type a : directly involved in charge transfer steps type b : not involved in charge transfer steps TABLE 2.-FORMATION OF CLASS 1 INTERMEDIATES examples of intermediates in homogeneous solution + e- l.preceding reactions : A f I1 f I, + P- 2. consecutive reactions : A Y+ 11 f 1; + P- 3. multiple charge transfer : A+ + I1 + I2 + P- 4. disproportionation : A+ + I ; 21 + A++ P- + e- -t e- + e- -t e- There are intermediate cases where this distinction is rather questionable. E.g., H-atoms as generated in electrode reactions will normally belong to class 3 of this scheme. However, when H-atoms are generated by photoemission of electrons, as discussed in the papers of Barker, Pleskov and Schiffrin, they appear also in solution. Furthermore, it is possible and even of frequent occurrence that intermediates of different types are found in one particular electrode reaction. In spite of these ambi- guities, I hope that this classification can be helpful for the following discussion.With regard to the electrochemical reaction, there is another important aspect which has serious consequences for the detection of intermediates. This is the question whether a particular intermediate is involved directly in the transfer of electrical * The literature quoted in an introductory paper can provide only a few examples and i s necessarily far from complete.H. GERISCHER 9 charge through the interface or not. Accordingly, two types of intermediates (a and b) are distinguished in table I . In the first case (type a) a much higher sensitivity of detection can be reached than in the latter one (type b). The following 3 tables (2,3 and 4) give some types of composite reactions in which these classes of intermediates are formed.TABLE 3.-FORMATION OF CLASS 2 INTERMEDIATES examples for chemisorbed intermediates (b) dissociative chemisorption : XY,,l f Xad+ Yad (c) chemisorptive discharge : x,+,~ +- Xad f: &,I ( d ) recombinative discharge : XZl+ Yad + XYad + XY,,l (a) chemisorption : xso1 * Xad + e- + e- 11 YSOl TABLE 4.-EXAMPLES OF CLASS 3 INTERMEDIATES intermediate states of the electrode surface 1. surface compounds on metals or semiconductors 2. electronic surface states on semiconductors 3. kink sites 4. two-dimensional nuclei 5. lattice defects in the surface Most of the discussion of the meeting will be concerned with intermediates of class 1 and 2. The intermediates listed in class 3 are found mainly during the formation o r dissolution of a solid electrode material.Their exact description is one of the most difficult tasks, which will be touched-on in the papers given by Fleischmann, by McIntyre and by Kolb. METHODS FOR DETECTING INTERMEDIATES I think it is useful to distinguish between indirect and direct methods of detecting intermediates. By indirect methods I mean the classical analysis of stationary (current, voltage) curves as well as the interpretation of non-stationary experiments. Direct methods are all techniques where the products are quickly transferred from the electrode-solution interface to an analyzing system and in situ spectroscopy. 1. INDIRECT METHODS (a) The analysis of steady state current voltage curves can give the rate constants of the charge transfer step provided inass transfer plays a minor role in a sufficiently wide range of the (current, voltage) curve or the mass transfer influence can be taken into account correctly (e.g., Tafel plots corrected for concentration polarization). Characteristic for this step is the exchange current density io at equilibrium. When intermediates of preceding or consecutive reactions are directly involved in the charge transfer step (intermediates of type a), the dependence of io on the reactant concentra- tions can give the reaction order and in this way evidence for such intermediate^.^-^10 INTERMEDIATES I N ELECTROCHEMICAL REACTIONS Although the results are not always unambiguous, the conclusions are quite convincing in many systems.Slow preceding reactions lead to reaction-controlled limiting currents.From the concentration dependence of such limiting currents one can similarly get conclusions on the intermediates formed in such a reaction.6 However, the steady state does not indicate intermediates formed in irreversible consecutive reactions after the charge transfer step, because such products cannot influence the (current, voltage) behaviour. Highly sophisticated calculations have been made for polarographic (current, voltage) curves beginning with the work in the Prague school of polarography and later extending over many other places in the world.8* Thanks to the excellent reproducibility of polarographic experiments, the sensitivity of the polarographic (current, voltage) curves for distinguishing different reaction mechanisms is optimal.The situation is much more difficult with solid electrodes. In all cases where intermediates of type 2 or 3 are formed, the enormous variability of the interaction between surface sites and chemisorbed species and its dependence on structural factors give an uncertainty in the interpretation of (current, voltage) curves which does not permit analysis comparable in exactness to the polarographic case. This is often overlooked in mechanistic interpretations of (current, voltage) curves, when straightforward applications of formal react ion kinetics are used in surface reactions of solid electrodes. In summary we can conclude that steady state (current, voltage) curves will often indicate that the mechanism is complex and that intermediates are involved. The identification of the species, however, is possible only in favourable cases.(b) More powerful are nonstationary or relaxation experiments. The reason is that transport is much faster in the initial state of a sharply commenced electrode reaction, as by a step function, or during rapid recycling procedures. The use of this idea goes back to Le Blanc’s experiments with alternating current electrolysis.l* The first application of linear relaxation techniques was made by D o h and Ershler l * and by Randles l 3 who analyzed the frequency dependence of electrode impedances. This was discussed during a predecessor of this conference-the Faraday Discussion on Electrode Reactions in 1947-which marked the beginning of an extremely success- ful period in the development of electrochemical kinetics.Since then, a great number of different nonstationary techniques have been developed and many are still under investigation, as this conference shows in the papers contributed by Albery, Armstrong, Conway, Eggins, Epelboin, Fleischmann, Harrison and their coworkers. It might be interesting to ask why so many different relaxation methods have been developed in electrochemistry compared with the few, but extremely powerful relaxa- tion methods, for homogeneous reactions in solution. The reason is that the kinetic equations of the processes involved in heterogeneous reactions, even for small devi- ations from equilibrium cannot be transformed into a finite set of independent linear relaxation equations with a corresponding number of relaxation times, because diffusion cannot be described by one single relaxation time.Characteristic time constants for electrode processes only denote a relation between diffusion and com- petitive reactions as rate controlling steps. For example, the time constant in a potentiostatic relaxation experiment refers to the transition from charge transfer con- trol to diffusion contr01.l~ Relaxation experiments with electrode reactions are like rate experiments in open systems which, in the main, cannot be described by one set of relaxation tirnes.lS Non-stationary experiments with electrode reactions have, therefore, to be analyzed by solving the full set of differential equations with the appropriate boundary con- ditions. The advantage over homogeneous reaction kinetics, where one faces similarH .GERISCHER 11 problems in the case of nonlinear rate equations, is, however, that one of the observ- ables, current or overvoltage, can be externally controlled with great accuracy. This allows an adjustment of the external conditions of the experiments to give an optimal information for particular reaction steps and is the reason why non-stationary electrochemical experiments have been developed in such a variety of techniques. I wish to make only a few remarks on the sensitivity of non-stationary techniques for detecting intermediates. As already mentioned, the sensitivity of all indirect methods is highest for intermediates which take part directly in the charge transfer reaction. Impedance measurements at equilibrium are most sensitive for measuring exchange current densities, though the double layer effects have to be evaluated.16* l 7 The method where the rate of the double layer capacity discharge is used for the kinetic analysis, the so-called coulostatic as applied in the paper of Barker and coworkers in this discussion, seems to be the fastest linear relaxation technique available at the present time.Intermediates of type 2 (chemisorbed) which are formed by a charge transfer step can often be detected by pseudo-capacity This will be discussed later in the papers of Armstrong and of Epelboin. The highest sensitivity is reached here if the potential-dependent chemisorption reaction is fast. The limitation comes from the uncertainty of the double layer capacity which has to be subtracted.In favourable cases, less than 1 % of a monolayer can be detected in the capacitive part of the impedance. If the formation of the adsorbed intermediate is slow, the detection is much less sensitive because other Faradaic currents interfere then with the pseudo- capacitive contribution. In some cases the intermediates formed in the surface can inhibit or catalyze the overall electrode reaction, when a very high sensitivity is observed. Catalytic effects can cause even inductive components in the electrode impedance,24 as will be further discussed in the contribution of Epelboin and coworkers. For detecting intermediates formed in consecutive reactions it is best to invert the sign of the current momentarily. Such relaxation techniques with a variety of step functions or ramps for current or potential are very p o w e r f ~ l .~ ~ - ~ ~ The lifetime of the intermediates must be long enough to survive the time needed to charge the double layer for reversing the current, which takes at least a few micro seconds. I n spite of the technical progress in instrumentation and computerized evaluation of the results, uncertainty about double layer effects in electrode kinetics still presents a serious limitation to the analysis of fast reactions and the detection of intermediates in small concentration. Even the most exact measurements often cannot give an unambiguous picture of what intermediates are involved. Such a proof needs direct evidence from other methods. 2. DIRECT METHODS A great many of the most recent developments haveconcentrated on direct methods.1 . CLASSICAL ANALYSIS BY SAMPLE EXTRACTION The classical direct method of chemical kinetics for detecting intermediates is to take a sample from the reacting solution and to transfer it to the analyzing system. The paper of Bruckenstein and coworkers gives an example of this technique. For electrode reactions it is useful to take the sample from the close vicinity ofthe electrode. One can try to stop further progress of consecutive reactions by rapid cooling or by chemical means of reaction interruption. Such techniques have been rather in- frequently used in electrochemistry. They are applicable for intermediates with long12 INTERMEDIATES I N ELECTROCHEMICAL REACTIONS enough life times for the analysis to be done outside the electrolytic cell.Much shorter life times are feasible if the analyzing system is located inside the cell very close to the generating electrode. An example of this technique is given in the paper of Kastening at this Discussion. Transmission spectroscopy with transfer of reaction products via a flow channel from the electrode to a window in the immediate vicinity has been described by Janata and Mark.28 Electrochemists often prefer electrochemical analyzing techniques which can be applied without taking a sample from the solution. By means of an auxiliary electrode which can be biased independently of the working electrode and located in its close vicinity, a coulometric analysis of intermediates is obtained if they can be reduced or oxidized electrochemically.The most elegant technique of this kind is the rotating ring-disc electrode of Frumkin and Nekrasov2g with the theory worked out by Levi~h.~O We shall hear about applications of this technique later on in papers of Nekrasov and of Albery. Similar arrangements have been constructed with rotating ring-ring electrode^,^^ or with the electrodes in the wall of a streaming channel where the distance between the generator and analyzing electrode can be made extremely Such techniques can obviously only be applied to intermediates of class 1, existing in solution, and in a more limited way to class 2 intermediates, if they can be desorbed partially. 2. IN SITU SPECTROSCOPIC TECHNIQUES The most exact direct method of detecting intermediates is in situ spectroscopy.All wavelength ranges and all techniques of spectroscopy are applicable in principle. However, there are limitations which depend on the properties of the solution and of the electrode. At least one of these or both have to be transparent for the incoming and outgoing electromagnetic signal. Furthermore, a wavelength has to be found where the expected intermediates differ sufficiently in absorbance from the reactants and the products. In this regard, intermediates in organic electrode reactions are very favourable since they often have the character of radicals. Then, th.eir absorp- tion edge in the u.-v. or visible range is usually shifted to a longer wavelength com- pared with the parent compounds. Their paramagnetism can be used for e.s.r. measurements. Spectroscopy can also be made especially sensitive by modulation technique^.^^ For many intermediates which are generated in low concentration this is the only way of obtaining spectroscopic evidence.In electrode reactions, we are in the for- tunate position of being able to monitor the rate of the reaction externally by current or potential control. These techniques are rather new and still undergoing rapid development. A great deal of discussion on them will take place at this meeting. (a) TRANSMISSION SPECTROSCOPY This technique has been mostly used with optically transparent electrodes, made either by using highly doped layers of semiconductors (with a large enough band gap to be transparent, like SnO,) or using semitransparent evaporated layers of noble metals on glass or quartz plates.The reflected light beam from a smooth metal surface can also be used. Since the reflectivity of the metal surface depends on surface preparation, wave number and electrode potential, this has to be taken into account in the latter case. Examples of these techniques will be given in the papers of Kuwana and of Bewick and their coworkers. An excellent review has just been given by K ~ w a n a . ~ ~ Since the whole content of the solution is transmitted by the light beam, thisH. GERISCHER 13 technique is most sensitive for intermediates diffusing rapidly into the solution and being generated at high current density. Modulation of the generation by varying the current densities increases the sensitivity by orders of magnitude and allows signal averaging techniques to be applied.As Kuwana has pointed out in his con- tribution to this Discussion, intermediates which are present in concentrations of M within a diffusion layer of about cm thickness can be detected if their absorbance is high enough. Problems caused by the absorbance of the solution can be seriously reduced by using thin layer cells for optical transmission, as developed by Murrey et ~ 1 1 . ~ ~ (b) ATTENUATED TOTAL REFLECTION SPECTROSCOPY Since the highest concentration of intermediates is normally to be expected at the interface or in the solution close to the surface, a light beam with internal total reflection at the electrode surface is especially powerful. It has an especial sensitivity for adsorbed or chemisorbed species.An excellent review has just been published by H a n ~ e n . ~ ~ The totally reflected light beam " sees '' from the solution only the depth of about 2/10 next to the interface. The formation of new surface compounds can be detected if their spectrum differs sufficiently from that of the other components. Adsorbed species of high absorbance (like dyes) have been detected down to 1 % of a monolayer. This method is most suitable for intermediates of class 2 and 3. The high sensitivity causes also some problems. The modulation of the rate of the electrode reaction via potential or charge also modulates the optical properties of the electrode surface and of the double layer. These effects can interfere with the spectroscopic analysis of intermediates. (C) SPECULAR REFLECTION SPECTROSCOPY The reflectivity at the surface of a solid is also very sensitive to chemical changes on the interface.This is most pronounced in the visible and u.-v. region where even metals lose their extreme reflectivity and cease to shield the electrical field of an electromagnetic wave completely. The result is that the electric field has no node in the surface (as for metals in the i.-r. range) and, therefore, interacts strongly with chemisorbed or any other species in the interface. The electrical field in the surface depends on the plane of polarization and the angle of incidence, which increases the obtainable information in a very specific way. An authoritative review has recently been published by M ~ I n t y r e , ~ ~ who will also give a paper on this subject here.It has been claimed that 1 % of a monolayer can be detected in this The interpretation of the results of such measurements is still not unambiguous, since the reflectivity variations are produced by a combination of several effects. The electrical field influences the optical constants of the electrode as well as the forma- tion of new surface compounds or a change in double layer composition. Recent studies have been devoted to the optical properties of monolayers and submonolayers of cathodically deposited metal atoms on solid substrates which can give insight into the electronic interaction between such atoms mutually and with the substrate.38 This problem will be discussed in the paper of Kolb. (d) ELECTRON SPIN RESONANCE In spite of the high absolute sensitivity of e.s.r. measurements, the in situ applica- tion of this technique for electrode reactions 39* 40 has proved to be more limited than originally supposed.The reason is that the resonator for microwaves imposes14 INTERMEDIATES IN ELECTROCHEMICAL REACTIONS serious limitations on the construction of the electrolytic cell with the result that only relatively small current densities can be applied. The high reactivity of radicals, especially resulting in disproportionation reactions, prohibits the build-up of a high concentration. The advantage of e.s.r. measurements is in its great power of radical identification if high resolution spectra can be obtained. This, however, needs a much higher concentration than necessary for the mere indication of a radical.Examples of the application of these techniques will be discussed at this meeting by Kastening and by Dohrmann. A summarizing survey of the direct methods, giving their field of application and their approximate sensitivity, with rather optimistic assumptions, is shown in table 5. TABLE 5.-DIRECT METHODS FOR DETECTING INTERMEDIATES method classical analysis by probe taking application sensitivity, in terms 01‘ to intermediates 1st order decay times class 1 1 - lo-’ s electroanalysis in situ class 1 10-3 s transmission spectroscopy class 1 (class 2) 10-4 s attenuated internal class 2 and 3 10-5- 10-6 s specul ar reflection class 3, 2 10-5- 10-6 s e.s.r. class 1 10-2- 10-3 s reflection spectroscopy (class 1) spectroscopy and 1 We see that the direct methods can even detect very short lived intermediates.However, we must not forget that these figures are calculated for very favourable cases. All spectroscopic methods decrease in sensitivity directly with decreasing difference in molar absorbance. The analyzing electrochemical methods suffer often from interference with other reactions. Therefore, in many systems only a much lower sensitivity can be obtained, and it remains the task of the experimentalist to select the optimal method for this particular system. THE PROSPECT FOR OUR SUBJECT This Discussion promises to give an excellent representation of the present state and will certainly outline the trends of future developments. I think it needs not much prophetic skill to predict that combinations of various techniques, direct and indirect ones, will be most successful in this field as is already indicated in some of the papers presented here.I am not so pessimistic for the future of this kind of re- search as Tafel was, when he published his ideas on the hydrogen evolution reaction. He wrote in his paper of 1904 the following sentences, which I quote here in transla- tion: “ The problem of electrode polarisation in electrolysis has been studied scienti- fically for about one hundred years. It is, therefore, scarcely possible to find some new aspects which previously have not been already touched either in experiments or in speculations.” I am convinced that also in a field with such a long tradition, as electrode reactions have, new aspects and surely new intermediates will be found in the future.In the spirit of the following Discussion, I quote a piece of advice, given by Sir Francis Bacon, than which nothing can better guarantee the success of such a meeting.H . GERISCHER 15 ’‘ He that questioneth iiiuch shall learn much, and content much ; but especially if he apply his questions to the skill of the persons whom he asketh; for he shall give them occasion to please themselves in speaking, and hiinself shall continually gather knowledge. But let his questions not be troublesome, for that is fit for a poser ; and let him be sure to leave other nien their turns to speak.” J. Tafel, 2. phys. Chenz., 1900, 34, 199. J. Tafel, Z. phys. Chem., 1904, 50, 641. K. J. Vetter, 2. phys. Chem., 1950, 194, 284; 2.Elektrochem., 1951, 55, 121. R. Parsons, Trans. Faraday SOC., 1951,47, 1332. H. Gerischer, 2. Elektrochem., 1953, 57, 604 ; Z. phys. Chetn., 1953, 202, 292. H. Gerischer and K. J. Vetter, Z. phys. Chem., 1951, 197, 92. R. BrdiEka and K. Wiesner, Naturwiss., 1943,31,247 ; Coll. Czech. Chem. Comm., 1947,12,39. R. BrdiEka, V. HannS and J. Kouteck? in Progress in Polarography, Vol. 1, ed. P. Zuman and 1. M. Kolthoff (Interscience Publ., New York, 1962), p. 145. R. Tamamushi and P. Sato in Progress in Polarography, Vol. 3, ed. P. Zuman, L. Meites and 1. M. Kolthoff (Interscience Publ., New York, 1972), p. 1. l o M. LeBlanc and K. Schick, Z. phys. Chem., 1903, 46, 213 ; 2. Elektrocliem., 1903, 9, 636. l 1 P. Dolin and B. Ershler, Acta Physicochim. U.R.S.S., 1940, 13, 747.l 2 B. Ershler, Disc. Faraday Sac., 1947, 1, 269. l4 H. Gerischer and W. Vielstich, 2. phys. Chem. (Frankfurt), 1955, 3, 16. J. E. B. Randles, Disc. Faraday SOC., 1947, 1, 11. M. Eigen and L. De Mayer in Techniques of Organic Chemistry, VIII, 2, ed. A. Weissberger (Interscience, New York, 1963), p. 895. A. N. Frumkin, Z. phys. Chem., 1933,164,121. R. Parsons in Adv. Electrochem. Electrochem. Eng., Vol. 1, ed. P. Delahay (Interscience Publ. New York, 1961), p. 1. l 8 G. C. Barker, Trais. Symp. Electrode Processes, ed. E. Yeager (J. Wiley, New York, 1961), D. 375. l9 P. Delahay, J. Phys. Chem., 1962,66,2204. 2o W. H. Reinmuth and C. E. Wilson, Anal. Chem., 1962,34, 1159. 21 H. Gerischer, 2. phys. Chem., 1951, 198,286. 22 W. Lorenz, Z. Elektrochem., 1954,58,912. 23 H. A. Laitinen and J. E. B. Randles, Trans. Faraday SOC., 1955, 51, 54. 24 H. Gerischer and W. Mehl, 2. Elektrochem., 1955, 59, 1049. 2 5 M. Kalousek, Coll. Czech. Comm., 1948, 13, 105. 27 D. G. Davis, Adv. Electroanal. Chern., 1966, 1, 157. 28 J. Janata and H. B. Mark, Jr., Anal. Chem., 1967,39,1896. 29 A. N. Frumkin and L. N. Nekrasov, Dokl. Akad. Nauk S.S.S.K., 1959, 126, 115. 30 A. N. Frumkin, L. Nekrasov, B. Levich and Yu. Ivanov, J. Efectroanal. Chem., 1959, 1, 84. 31 K. E. Heusler and H. Schurig, 2. phys. Chem. (Frankfurt), 1965, 47, 117. 32 H. Gerischer, I. Mattes and R. Braun, J. Electroanal. Chem., 1965, 10, 553. 33 M. Cardona, Modulation Spectroscopy (Academic Press, New York, 1969). 34 T. Kuwana, Ber. Bunsenges. phys. Chem., 1973,77, 35 R. W. Murrey, W. R. Heinemann and G. W. O’Dorn, Anul. Chem., 1967,39, 1666. 36 W. N. Hansen in Adv. Electrochem. Electrochem. Eng., 1973, 9, 1. 37 J. D. E. McIntyre in Ado. Electrochem. Electrochem. Eng., 1973, 9, 61. 38 D. M. Kolb, Ber. Bunsenges. phys. Chem., 1973, 77, 39 D. H. Geske and A. H. Maki, J. Amer. Chem. SOC., 1960,82,2671. 40 R. Koopmann and H. Gerischer, Ber. Bunsenges. phys. Chetn., 1966, 70, 118, 127. W. Jaenicke and H. Hoffmann, Ber. Bunsenges. phys. Chem., 1962, 66,803, 814.

ISSN:0301-7249    

DOI:10.1039/DC9735600007

出版商:RSC    

年代:1973

数据来源: RSC

摘要:

Spectro-electrochemical Methods in the Study of Short-lived Intermediates BY ROBERT F. BROMAN, WILLIAM R. HEINEMAN: AND THEODORE KUWANA * Department of Chemistry, Ohio State University, Columbus, Ohio 43210, U.S.A. Received 17th July, 1973 The theoretical and experimental state of the art of spectro-electrochemistry as directed toward the study of short-lived intermediates is discussed. ParticuIar emphasis is placed on the detection, characterization and kinetic monitoring of reactive intermediates during potential puIse and relaxation periods. Nucleophilic addition to radical dations with digital simulation of the possible mechanisms for the pulse-relaxation are presented using the thianthrene system as a working model. There has been a recent accelerated growth in the number of investigators who have utilized some optical method to assist in the elucidation of electrode processes.This growth has been stimulated by the potentialities of obtaining fundamental information heretofore unattainable with other methods, by the multitude of possible application areas and especially, as from the beginning, by the tremendous technical developments in the optical field. These optically coupled electrochemical experiments, which we shall refer to as spectro-electrochemistry, are presently moving in two directions. The first is primarily the study of the mechanistic aspects of electrode processes with emphasis on the detection, characterization and kinetic monitoring of electrogenerated short-lived species, and utilizes in most instances an optically transparent electrode (ote).The principal optical methods have been normal transmission spectroscopy or internal reflection spectroscopy (IRS), and they provide a convenient and specific in situ probe to follow some species to assist in the interpretation of an electrode mechanism. The second direction is the optical study of phenomena associated with the electrode surface extending from the " interphases " of the surface to the diffused double layer region.2 Those in this latter group have employed principally reflection methods, i.e., ellipsometry and specula reflection, from solid surfaces (or liquid in the case of mercury). Internal reflection spectroscopy has also played a significant role in these surface studies with thin conductor film electrodes.Possibly the most significant experimental achievement during the last few years has been the development of instrumentation capable of detecting and evaluating extremely small changes in the level of light. For example, changes in the normalized reflectance (ARIR) on the order of 2-5 parts in lo6 have been reported for the electro- modulation (reflectance) technique. And, with the use of signal averaging instru- mentation, changes in optical absorbance (AA) of a few parts in lo6 can be achieved for the kinetic monitoring of electrogenerated species.l Thus, it has become possible optically to monitor minute changes that occur as a consequence of some electro- chemical perturbation. These changes have been ascribed to less-than-monolayer coverages of surfaces by adsorbed materials, to changes in surface roughness, to the permanent address :-Department of Chemistry, University of Cincinnati, Cincinnati, Ohio.16R. F. BROMAN, W. R. HEINEMAN AND T. KUWANA 17 electromodulation of inner band transitions and surface charge density, and to changes in the structure or ionic concentration at the electrode-solution interface. With respect to the study of electrogenerated short-lived intermediates, the high sensitivity at short times allows kinetic evaluation of bimolecular rates up to the diffusional limit. In this paper, the discussion will be on the theoretical and experimental state of the art of spectro-electrochemistry with particular emphasis placed on the detection, characterization and kinetic monitoring of reactive intermediates during potential pulse and relaxation periods.EXPERIMENTAL Electrochemical and optical instrumentation for potential step with concurrent spectral monitoring through an optically transparent electrode by rapid scanning and signal averaging have been rep~rted.~-~ In the thianthrene studies a glass electrochemical cell similar to the Lucite cell described by Hawkridge and Kuwana was used with an internal coil of platinum wire serving as the auxiliary electrode. The reference electrode consisted of a Ag/AgCl wire in 1 F aqueous NaCI. Solution contact was made through an asbestos-fibre junction in soft glass and a Luggin capillary constructed of Teflon tubing and filled with supporting electrolyte and acetonitrile. The optically transparent electrode was platinum on quartz with a resistance of ca.20 ohm sq-l. Baker A.R. grade acetonitrile was dried by addition of neutral alumina which was previously dried at 450°C for 24 h. Thianthrene and tetraethylammonium perchlorate (TEAP) were recrystallized. The latter was dried under vacuum at 70°C for 12 h ; 2.0 mM solutions of thianthrene in 0.5 F TEAP in acetonitrile were prepared and stored under dry nitrogen. Water was added to a known volume of the soIution in the electrochemical cell with a microlitre syringe. GENERAL CONSIDERATIONS Transmission spectrophotometry is commonly used because of its simplicity. In transmission (fig. la) the light beam traverses the entire cell and the beam is atten- uated by whatever light interaction changes occur as initiated by the electrochemical perturbation.Transmission methods fulfil most of the usual needs. For IRS (fig. lb) multiple reflection can improve sensitivity to an intermediate in the solution phase, but only up to a point. Scattering and polarization effects for each reflection must be considered. If at all possible, it is best to use only one reflection for quantitative work. Honiogeneity of the surface is also a problem when one uses several reflections. The light beam penetrates into the solution phase; the depth of penetration, 6, is dependent on the wavelength, the incident angle, 6, and the complex refractive indices of each phase involved. Thus, the optical cell width in the solution phase can be varied and controlled, and this depth can be easily calculated.8 Not only can signal averaging improve the signal-to-noise ratio but the optical coupling at the electrode-solution interface can be enhanced if the thin conductor film is some multiple of the wavelength used,6 i.e., similar to a quarter-wave.The conductor phase then acts as an optical amplifier and a gain of several fold can be achieved by adjustment of optical and conductor phase parameters. The initial problem in spectroelectrochemistry is quantitatively to describe the optical absorbance ( A ) as a function of time during the application of a given electro- chemical method. Because of the finite surface resistance with the concomitant problem of maintaining an equipotential surface and the difficulty of numerically solving diffusion-kinetic equations, the electrochemical method of choice has been chronoamperometry (potential-step).The boundary value problems and the18 SPECTRO-ELECTROCHEMICAL METHODS - glass s o l ~ t io n -- conducting f i l m a b FIG. 1.-Spectro-electrochemistry : a. transmission, A = -log ] / I o = E Cx,r dx, d = cell thick- ness; b. IRS, A = N,fp CX,* exp(-xx/6)dx, 6 = penetration depth of beam, A = complex refractive index, H = thickness of conducting film, Neff = sensitivity factor dependent on film and optical parameters. J: 1: A 6 0.0008r P 0.0002 0 0 1.0 2.0 3.0 4.0 t/ms FIG. 2.-" ec mechanism ", A +A* Oe-, A +Z+P ; digital simulations. A. Concentration- distance profiles for A(- -), A? (-), Z(. * *) and P(* - .) ; (a) 0 ms, (b) 1 ms and (4 4 ms ; kl = 10' 1. mol-1 s-1 ; c i = ci = 1 mM, C i * - = Cp = 0 ; D = cm2 s-'.B. Absorbance-time curves for kl values of (a) 0, (6) lo5, (c) lo6 and (d) lo7 1. mol-' s-' ; monitor A&- optically, E = lo4 1. mol-' cm-'. k iR . F. BROMAN, W . R . HEINEMAN AND T . K U W A N A 19 subsequent solutions have been numerically evaluated by the computer-digital simulation method of Feldberg.9 Briefly, in the case of a simple electron transfer reaction, A+ A * a & e- (1) the formation of A? (oxidation) or A? (reduction) at a diffusion controlled rate (chronoamperometry) results in the relationship. L A(t) = -EC'JE Jn where E is the molar absorptivity of the monitored species, A** ; D is the diffusion coefficient of species A in cm2 s-' ; and c" is the bulk concentration of A in moll.-'. In this case, absorbance is a linear function of ,/t.For an " ec mechanism " A+A** e- k l A**+Z+B the diffusion-kinetic equation to be evaluated is a[A'.] d'[A*.] - D F - k,[A**][Z] at (4) and the C-x-t profiles computed by digital simulation for kl = lo7 1. mol-l s-l, E A i . = 1.0 x lo4 1. mol-1 cm-l, and DA = 4.0 x cm2 s-l are given in fig. 2A. The thickness of the diffusion layer moves as a function of Jt. The ( A , t ) response for various values of kl for this mechanism under second order conditions are shown in fig. 2B where A*- is monitored optically. Reaction (3) is written as an irreversible one, however, a Kes for the reaction incorporating kel for the back reaction can be included in the simulation. The magnitude of Keg then governs whether the back reaction can be ignored or not.For the condition [Z] 4 [A], the pseudo 1st order rate constant, k', is k' = kl[Z] ( 5 ) and as shown by Grant,'O the A response goes through a maximum during a chrono- absorptometric experiment. The time at which the absorbance reaches a maximum value, t,,,, is related to k' by and should agree with the k' determined from open circuit " relaxation " which is the conventional 1st order kinetic decay (t+ = 0.6931k'). From the plot of (A, Jt), the maximum absorbance value at t,,, is easily evaluated for various k' values, and usually, the E JD values can be evaluated from the initial linear portion at short times. The experimental parameters and conditions are summarized in fig. 3 where the logarithmic values of A , x and t are plotted. The diffusion layer thickness, x, is approximated as x = 2 JDf and is the Nernst reaction layer thickness.At 1 ins, it can be seen that the reaction layer thickness is the order of 10-4cm and the A is between and units. The sensitivity of most commercial spectrophoto- meters is about For set wavelength monitoring of a reactive intermediate at times less than 1 s, it is necessary to devise one's own optical apparatus. For times approaching to s, the t,,, = 0.85/k' (6) unit with a response time usually greater than 1 s.20 SPECTRO-ELECTROCHEMICAL METHODS reaction layer is now from several hundred to a thousand Angstroms thick and the required optical sensitivity is less than 1 part in 10 000. To achieve optical monitoring of A** at these times, modulation coupled with signal averaging instrumentation is required.log(tls) FIG. 3.-Relationship of experimental parameters for kinetic measurements by spectro-electro- chemistry ; E = 1.0 x lo4 1. moI-' cm-', D = 4.0 x cm2 s-' ; C" = I .O x M. For the pseudo 1st order " ec mechanism", the time axis can now be considered as t,,, in fig. 3, the time to reach the maximum absorbance during chronoabsorpto- metry and log k' is plotted on the top axis of fig. 3. There is almost a 1 : 1 reciprocal correspondence between the value of k' and t,,,. At tmax = s, the k' value is 0.85 x lo3 s-l. One can also approximate what time resolution is necessary for evaluation of bimolecular rates. s, kbi of lo8 1. mol s-' should be attainable using Z in the to M range. In our laboratory, times on the order of 25 ps for transmission signal averaging and 10-15 ps for IRS have been achieved.This enables kbi equal to about lo8 1. mol-' s-I for transmission and 3 x lo9 1. mol-1 s-l for IRS to be evaluated.'l Winograd l 2 has recently reported 3-5 p s resolution for signal averaged IRS using coulostatic charge- injection coupled with a potentiostat. Spectro-electrochemistry is thus capable of resolving kinetic rates very near to the bimolecular diffusional limit. For electrode mechanisms where species A** can be stabilized sufficiently, i.e., no 2 present, that an absorption spectrum can be obtained for the purpose of selecting a wavelength at which A*. absorbs, then it can be kinetically monitored at this known wavelength under conditions where it is reactive. However, if A*- reacts with solvent, dimerizes or is intrinsically unstable, spectral characterization by acquisition of spectra during electrogeneration may be difficult.Such acquisitions have been performed in our laboratory with use of a rapid scanning spectrophotometer (RSS) which employs an oscillating mirror to sweep the dispersed spectrum across the exit slit. Two such spectrophotometers are presently operational and provide spectral access to wavelength range of 2400 to 9500A with a variable scan rate up to a maxi- mum of 1 ms per spectrum at a wavelength width of 2500A. The absorbance sensi- tivity at a signal to noise of unity is about 0.005 unit with a wavelength resolution of 5-10 A, depending on the wavelength range. Spectral acquisition by use of magnetic tape recording has been previously discussed.Computer controlled acquisition, signal averaging and analysis are being developed using a Nova 800 minicomputer. Similar efforts are underway in at least three other 1ab0ratories.l~ For example at t =R . F . BROMAN, W. R . HEINEMAN A N D T. KUWANA 21 PULSE-RELAXATION KINETICS FOR THE STUDY OF HOMOGENEOUS REACTIONS Spectro-electrochemistry provides the unique opportunity of following an electro- generated species both during an applied electrochemical perturbation and then during a subsequent open-circuit relaxation period, independent of the electrochem- istry. The mechanism derived from these two time periods should be consistent and should assist in the delineation of any steps in the mechanism which involve secondary or post-heterogeneous electron transfer.A good example of the utility of pulse-relaxation is given by the study of the oxidation of DPA in the presence of water in acetonitrile by Grant.'O Grant found that the kinetic results were not in agreement from analysis of the (A, t ) periods be- tween pulse and relaxation when DPA+- was being optically monitored. Blount '' re-examined the results and concluded that the previous analysis did not properly account for the reaction in which DPA was regenerated (reaction (9)). This reaction provides a higher flux of DPA to the electrode surface and consequently, a higher concentration of DPA+* than expected from reactions (7) and (8) alone. This must be taken into consideration during the time the potential is applied.When this is done, the kinetic results between pulse and relaxation are consistent with the half- regeneration mechanism of Sioda,16 namely DPA+DPA+. + e- (7) DPA+*+ H,O+DPA(OH)*+ H+ (8) (9) DPA(OH)++ H20-+DPA(OH)2 +H+. (10) ki k2 k3 DPA(0H)- + DPA+-+DPA(OH)+ + DPA In the analysis, it was assumed that k2%k1, and that steady state kinetic conditions could be imposed on the concentration of DPA(OH)*. As such, the rate of disappear- ance of DPA+* during open circuiting was given by d[DPA+*]/dt = - 2k',[DPA+*] where k', = kl[HZO]. This mechanism has been recently re-examined using pyridine instead of water.17 The (A, t ) behaviour for the addition of py to DPA+* under true second order condi- tions supported the half-regeneration mechanism. The experimental data and simulated curves for the normalized absorbance during potential pulse are shown in fig.4. Another example of the use of pulse-relaxation was in the evaluation of the rate constant for the reaction TAA+TAAf +e- (12) (13) k TAAf + CN-jTAA + CN* where the tri-p-anisylamine (TAA) is regenerated through reaction of the radical cation with cyanide The signal averaged pulse-relaxation pattern is shown in fig. 5. The pulse time was 2 0 0 p and the relaxation was followed for 3-5 ms. The normalized absorbance working curve was computed for the concentration ratios of C+AA to C&-and gave a value of k = 1.95( k0.18) x lo5 1. mol-l s-l which was in good agreement with rates obtained by spectro-electrochemistry using transmission single potential pulse, double potential pulse, and signal averaged IRS single potential pulse.22 SPECTRO-ELECTROCHEMICAL METHODS - -__ '."I r 0 .6 - 0.6 - R - P 0.4 - 0.2 1 A B C 0 - 1.0 0 I .o 2 .o Q J A p - l I - L - - i ~ I - - - l - _ - _ _ - -2.9 log(C,o,k10 FIG. 4.-Comparison of computer simulated spectro-electrochemical response for half-regeneration mechanism (solid lines) with experimentally determined values. (A) C p y / C ~ p ~ = 0.0514, kl = 1.69(f0.37) x lo4 M-I S-I ; (B) C p y / C ~ p ~ = 0.105, kl = 1.79(+0.44) x lo4 M-I s-' ; (C) CP~/CDPA = 0.207, kl = 1.78( f 0.33) x lo4 M-' S-' ; (D) C p y / C ~ p ~ = 0.484, k , = 2.13( & 0.42) x lo4 M-' S-' ; (E) cpy/cDpA = 1.022, kl = 1.84(*0.27)~ lo4 M-' S-'. Overall weighted average of kl = 1.86( k 0.38) x lo4 M-I s-' . Time span of data: 5 ms to 5 s ; temperature = 24.0f 0.5"C.(Reproduced with permission : H. N. Blount, J. Electroanal. Chem., 1973, 42,271). As pointed out previously, the limitation of this pulse-relaxation method is the rate of the open circuit electrode discharge relative to the rate of the homogeneous chemical reaction. The half-time of the homogeneous chemical reaction in question must be greater than the half-time of the electrode discharge. Also for any signal averaged experiments, the restoration of the initial solution conditions prior to each pulse is a necessary prerequisite. This is accomplished by the imposition of the initial working electrode potential (prior to pulse) and waiting for a time period which is 100 times the time of the applied potential pulse. Restoration at these time ratios is completed to more than 95 % of the original condition.I' Y I 1 I I I 1 1 , b'*ll 0 0.88 1.76 2.64 tlm FIG. 5.-Spectroelectrochemical open circuit relaxation for 0.5 mM TAA with 15 mM CN-, 200 p pulse, 800 repetitions for signal averaging.R . F . BROMAN, W. R . HEINEMAN A N D T. KUWANA 23 Pertinent to the discussion of nucleophilic addition to radical cations, the digital simulation of possible mechanisms for pulse-relaxation will be presented using the thianthrene system as a working model. Possible sequences of the reactions for thianthrene are similar to those of the previously discussed DPA system. The proposed mechanism for thianthrene in the presence of water are as follows : I. The " ece mechanism " ; Th-ThS-0 + e- (14) (1 5) ThOH-+ThOH+ + e- (16) ThOH++ThO + H+.(17) k i k - 1 Th+* + H,O+ThOH* + H+ 11. The " disproportionation " mechanism proposed by Shine and Murata 19* 2o - Th -+Th+- + e- k f Th+* + Th+-+Th + ThS2 kb k6 Th+2 + H20+Th0 +2H+. 111. And, the " half-regeneration mechanism " ; Th-+Th-k*+e- ki k - t k z ThOH. + Th+*+ThOH+ + Th k - 2 Th+* + H,O+ThOH* + Hf ThOH++ThO + H+. The diffusion-kinetic rate equations are for mechanism I, assuming k , >k-,, a[Th] - a2[Th] DT -- at for mechanism TI, assuming a pre-equilibrium for reaction (IS), and irreversibility of reaction (19), (25) a[Th] = $2CTh1 k6kf [Th+*12CH20] CThl 4- __- at ax2 kb24 SPECTRO-ELECTROCHEMICAL METHODS where k6 is written as and kf/kb = Keg and for mechanism 111, assuming again kl >k-, and kl <k,, so that a steady state approximation can be made with respect to the concentration of ThOH-, d[Th+*] a2[Th'*] = D - 2 k 1 [Th ' *] [ H, 01 at ax2 Results of digital simulations of the above mechanisms are expressed in terms of working curves in which the absorbance at any time t is normalized to the absorbance value at the moment the cell was open circuited (Aoc).The normalized value is plotted against the dimensionless parameter log(CH,okl, t). For certain mechanisms and experimental situations, a single working curve can be unique for all values of the rate constants and concentrations. The open circuit relaxation experiment was simulated by setting the fluxes of Th and Th+= equal to zero at the electrode surface after 500 time increments. Comparative working curves for normalized absorbance during open-circuit (relaxation) are shown in fig. 6.The shapes of the ece (mechanism I) and the half- regeneration (mechanism 111) curves are identical, but shifted by a factor of two because of the manner in which the rate is written (eqn (30)). (The chronoampero- metric and double potential-step electrochemical working curves have been published by Marcoux 21 for the ece and ec disproportionation mechanisms.) However, the shape of the curve for the disproportionation mechanism is different from those of the other two mechanisms. -20 -110 0 I .o 2 .o log(C;i,okd) FIG. 6.-Digitally simuIated spectro-electrochemical working curves for open circuit relaxation : A, disproportionation (mechanism 11) ; By ece (mechanism I) ; C, half-regeneration (mechanism 111).R .F . BROMAN, W. R . HEINEMAN AND T. KUWANA 25 In the simulation of mechanism I, the electrochemical current arises from the fluxes of both Th and ThOH- at the electrode surface. The simulation was stable, and 1000 time increments were sufficient for good definition of the normalized absorbance working curves. For the disproportionation mechanism, the stability of simulation was markedly poorer than the other mechanisms. It is believed, however, that the working curves drawn through the data points corresponding to the terminus of each simulation represent the solution for the system. Stability of the simulation for the half-regeneration mechanism was excellent and identical results were obtained for Chz0 = 20 C& and 40 C&.It should be pointed out that the omission of the possible electroactivity of the species ThOH. in mechanism 111 will not affect the open circuit results, particularly assuming that k , 9 kl and that the equilibrium constant for reaction (20) is large. However, it will be important in the simulation of the electro- chemical current flux and also during the chronoabsorptometry experiment (without open circuit) when Th+- is being optically monitored. PRELIMINARY STUDY OF THIANTHRENE+ WATER SYSTEM Mwata and Shine 1 9 9 2o have reported that thianthrene dication (Th") in aceto- nitrile undergoes a second-order reaction with water (mechanism 11). Results of their spectrophotometric studies on the decay of the absorbance of Th+= (&546nm = 8.5 x lo3 1.mol-' cm-I) in solutions containing Th, Th+.ClO,, and water indicated a 1.0 0.8 4 0 . 6 1 + 0.4 0.2 0 - 2.0 - 1.0 0 log(C1rzok 1 t 1 FIG. 7.-Comparison of simulated working curve for half-regeneration mechanism with experiment- ally determined values under conditions of open circuit relaxation, CH~O = 0.069 M, (0) 3.5 s pulse, (0) 2.0 s pulse. reaction of water with Th+2 which was produced by the rapid disproportionation of Thf-. A value of k6k,/k, of ca. 12 1. mol-' min-' was measured at 25°C. In a brief study, Parker and Eberson 22 presented evidence that in the electrochemical generation of the radical cation, the radical reacted with water, not Th+2. They26 SPECTRO-ELECTROCHEMICAL METHODS suggested the half-regeneration mechanism (mechanism 111) as being more tenable in view of their data.Another possible mechanism for consideration has been the so-called ece (mechanism I). This mechanism has been cited frequently in the liter- ature with respect to reactions of radical cations and anions with proton containing species. Another aspect of the " correct " electrochemical mechanism is provided by Marcoux.2 The cyclic voltammetry for the oxidation of Th in acetonitrile at a Pt-otc (quartz) gave two oxidative waves; the second was irreversible. The ( A , t ) behaviour for chronoabsorptometry during potential step deviated from the expected behaviour for all three mechanisms and more so for mechanism 11, the disproportionation. The pulse-relaxation data, however, followed closely mechanisms I and 111, and deviated markedly from mechanism 11.The experimental and calculated normalized absorb- ance against the log(CA2&lt) are plotted for mechanism I11 in fig. 7 for CG2* equal to 35 and 69 mM, respectively. Rate constants of 2.14 and 2.64 1. mol-1 s-l were obtained for pulses of 3.5 and 2.0 s respectively. The shape of the working curves (fig. 6) is very similar for mechanisms I and I11 since the kinetic pathways are similar in the absence of the electrode reaction. With the limited data presently available, it is not possible to decide between mechanisms I and ITI. The data, however, do not appear to support the disproportionation mechanism. There is a significant difference in concentration of the radical cation used here and in the experiments of Shine and Murata which may account for observable differences in the kinetic parameters.Perhaps the most likely mechanism, in view of the wide range of concentrations of species found in the electrode reaction layer, is one which involves a set of parallel reaction pathways, e.g., both disproportionation and half- regeneration as shown for radical anions of several aromatic hydrocarbons by Fujihara, et aZ.23 Attempts are underway to evaluate the kinetics of the reaction of Th+2 with water by stepping the potential to a value where ThS2 is the major product of the electron transfer reaction. Detailed spectroelectrochemical analysis under a wide range of reactant concentrations coupled with product analyses, is also in progress.24 The authors gratefully acknowledge the support of research by National Science W.R. H. acknowledges support by the Research Foundation Grant GP 31236. Corporation, Cottrell Grant 6720. N. Winograd and T. Kuwana, in Electroatialytical Chemistry, ed. A. J. Bard (Marcel Dekker, N.Y.), Vol. 7, in press. see papers-Optical Studies of Adsorbed Layers at Iiiterfaces, Symp. Faradav Soc., 1970, 4. A. Bewick and A. M. Tuxford, ref. (2), p. 114. J. W. Strojek, G. A. Gruver and T. Kuwana, Anal. Chetn., 1969, 41,481. H. N. Blount, N. Winograd and T. Kuwana, J . Phys. Chetn., 1970, 74, 3231. N. Winograd and T. Kuwana, Anal. Chem., 1971,43,252. F. M. Hawkridge and T. Kuwana, Anal. Chetn., 1973, 45, 1021. S. W. Feldberg, in Elecfruanulytical Chemistry, ed. A. J. Bard ((Marcel Dekker, N.Y.), Vol. 3. a W. N. Hansen, R. A. Osteryoung and T. Kuwana, Anal. Chem., 1966 38, 1810. l o G. C. Grant and T. Kuwana, J. Elecfroanal. Chem., 1970, 24, 11. l 1 N. Winograd and T. Kuwana, J. Amer. Chem. Soc., 1971,93,4343. l 2 J. E. Davis and N. Winograd, Anal. Chem., 1972, 44, 2152. G. A. Gruver and T. Kuwana, J. Electrounal. Chem., 1972, 36, 85. l4 private communications : R. W. Murray, University of North Carolina at Chapel Hill ; G. S. Wilson, University of Arizona ; H. B. Mark, University of Cincinnati ; (Rapid Scanning Spectrophotometer commercially available from Harrick Scientific, Ossining, N.Y., U.S. A.). l 5 H. N. Blount and T. Kuwana, J. Elecfroanal. Clwm., 1970, 27, 464. l6 J. E. Sioda, J. Phys. Chem., 1968, 72, 2322.R . F . BROMAN, W. R . HEINEMAN AND T . KUWANA l 7 H. N. Blount, J . Electroanal. Chem., 1973, 42, 271. " H. J. Shine and Y. Murata, J . Amer. Cheni. SOC., 1969, 91, 1872. 'O Y . Murata and H. J. Shine, J. Org. Chent., 1969, 34, 3368. '' V. D. Parker and L. Eberson, J. Amer. Chem. SOC., 1970,92,7488. 23 M. Fujihira, H. Suzuki and S. Hayano, J. Electroanal. Chem., 1971, 33, 393. 24 H. N. Blount, private communication. G. Manning, V. D. Parker and R. N. Adams, J . Anlet-. Cltem. SOC., 1969, 91, 4584. L. Marcoux, J . Amer. Chem. SOC., 1971, 93, 537. 27

ISSN:0301-7249    

DOI:10.1039/DC9735600016

出版商:RSC    

年代:1973

数据来源: RSC

摘要:

Photochemical Generation of Semiquinone Intermediates at a Rotating Semi-transparent Disc Electrode BY W. J. ALBERY, M. D. ARCHER," N. J. FIELD AND A. D. TURNER Physical Chemistry Laboratory, Oxford Received 1st June, 1973 Light from a mercury lamp shone through a semi-transparent electrode into a solution containing a mixture of anthraquinone and its corresponding hydroanthraquinone generates semiquinones inside the diffusion layer of the electrode. The semiquinone species can either be oxidized or reduced on the electrode. The rate constants for their electrochemical reactions and for their homogeneous recombination can be measured. At the potential of zero photocurrent the electrochemical rate con- stants for the reduction and oxidation of the semiquinone are equal. At this point the free energies for the first and second transition states in the electrochemical reduction of anthraquinone are equal.The variation of the potential of zero photocurrent with pH determines the electrochemical mechanism of the anthraquinone. The experiments to be described may be regarded as either the electrochemical detection of photochemical intermediates, or the photochemical generation of eIectro- chemical intermediates. In these experiments, light from a high-pressure mercury lamp is shone through a semi-transparent electrode. The light is absorbed by mole- cules in the solution. The advantage of using a semi-transparent electrode and shining the light through it, is that a relatively high percentage of the light reaching the solution (in our experiments -12 %) is absorbed in the diffusion layer close to the electrode.Thus, the photochemistry takes place within range of the electrode. We also rotate the electrode. This gives us control of the transport of species to and from the electrode and allows us to study the system in the steady state. The systems we have studied in most detail are 9,lO-anthraquinone 2,6-disulphonate and 9,lO- anthraquipone 1,5-disulphonate. These systems were chosen because work had been done both on the electrochemistry 1* and the photochemistry of quinone systems.3* It is convenient to describe our electrochemical reaction system using the scheme of squares ' 9 and the following notation, where A, B and C refer to the different oxidation states; the sulphonate groups are not shown.The labels En refer to transition states for electron transfer. 0 0- 0 A 0- Be- 0- C - * now at the Royal Institution, 21 Albemarle Street, London W1 28w. J . ALBERY, M. D. ARCHER, N . J . FIELD AND A. D . TURNER A +B7+ C= H+ 11 H + 11 H+ 11 29 e e El E2 e e 4 E4 AH+ +BH*+CH- H + 11 H + 11 H+ It e e AH:+ +BHf +CH, ES E6 We have studied the systems in moderately acid solutions (O<pH<6) and, except at the most basic end of the range, with respect to protonation the stable species are A, BH- and CH2. For the reduction of A to CH2 or the oxidation of CH2 to A a number of different routes through the scheme of squares can be envisaged. The particular route that is followed depends on the pH of the solution and the pK of the various transition states.6 In basic solutions, transition states near the top of the scheme of squares will be preferred; for more acid solutions, the route may shift further down the scheme and involve the more protonated transition states and intermediates near the bottom. In our buffered systems the proton transfers are all rapid and the rate-determining steps are the electron transfers.At reducing potentials, the rate-determining step involves a transition state on the left-hand side, while at oxidizing potentials it is on the right-hand side. In general at extremes of potential the first electron transfer along the route is rate determining6 The potential at which the rate-determining step shifts from the first to the second electron transfer or vice versa is not the same as the standard electrode potential or the potential at which no current flows.In this work we photochemically generate the BH- intermediates close to the electrode so that we can study the electrochemistry of BH. both as a function of electrode potential and as a function of pH. Such a study enables one to chart the routes through the scheme of squares, their change with pH, and to investigate the details of the rates of reaction of the intermediate semiquinone species. EXPERIMENTAL The apparatus is shown in fig. 1. The semi-transparent electrode consists of a Spectrosil quartz rod 4 m i diam. and 12 cm long which is coated on one end with a semi-transparent layer of Pt, by painting the end of the rod with Johnson & Matthey PBC 2532 Pt paint and firing at 1120 K.from the centre of the disc to the electrode contact and a transmittance in the u.-v. of 20 %. The rod is mounted in a hollow steel shaft which is rotated by a Velodyne motor and pulley system. Light from an Osram 250W ME/D high-pressure Hg lamp is focused on to the top of the rod, which acts as a waveguide. The circuit of fig. 2 is used to control the potential of the electrode. The circuit is first balanced with no light shining through the electrode by adjusting the potentiometer S2 until the summing point P is at earth potential and no signal is seen on Vp. The dark current is measured by the digital voltmeter VD. On opening the shutter the small difference in current caused by the light is amplified and measured on Vp; photo- currents as small as 1 nA can be measured in this way.The resulting layer has a resistance of less than 20 RESULTS AND DISCUSSION Fig. 3 and 4 show typical results for the current voltage curves of the dark and The following seven pieces of evidence show that the photochemical photocurrents. processes involved are :30 PHOTOCHEMI C A L GENERA T I ON I8 v A-+A* A* + CH2-+2BH* kn BH*+BH*+A+CHz. First, in order to observe a photocurrent one has to have both A and CH, (quinone and hydroquinone) present in the solution close to the electrode. In fig. 4D there was no CH2 present in the bulk of the solution, and a photocurrent is only observed at cathodic potentials when CH, is being generated by the dark current. Secondly, the photocurrent has in general a cathodic and an anodic branch, so that the product of the photochemical process can both be reduced and oxidized.The potential of zero photo-current varies with pH but is independent of the ratio of the bulk concen- trations of A and CH,. Thirdly, at potentials where the dark current for the reduction of A or the oxidation of CH2 becomes significant the photocurrent declines to zero. FK. 1 .-Diagram of apparatus. FIG. 2.-Operational amplifier circuit for measurement of photocurrents.w. J . ALBERY, M. D. ARCHER, N . J . FIELD A N D A . D . TURNER 31 EIV FIG. 3.--Current voltage curves for 9,lO-anthraquinone 2,6-disulphonate ; [A] = 22 niM, [CHJ = 8 mM, pH = 1.12. E is measured with respect to a saturated calomel electrode. Note the different scales for the photocurrent (0) and the dark current (A).pH 3.2 0.02 , E with FIG. 4.-Current voltage curves for 9,lO-anthraquinone 2,6-disulphonate at various p€ respect to S.C.E. Photocurrent, - ; dark current, - - -. In D there was no hydroquinone in the bulk of the solution. This is because when the overpotential is so large that A or CH2 react at the transport limited rate, the electrode fails to discriminate between the unexcited A or CH2 and the photoproduct BH.. For example, for the reduction, dark light 2 e hv 2 e 2 e electrode, all A-+CH2 solution, A+2BH* electrode, all 2BH-2CH2 all remaining A+CH232 PHOTOCHEMICAL GENERATION With the light there is a flux of BH. at the electrode but in the transport limited region this is exactly balanced by a drop in the flux of A. The same effect is found in the oxidation of CH2.Fourthly, the dependence of the size of the photocurrent on the intensity of the light I agrees with this photochemical scheme. We find that the dependence of i on I varies between I* and I t . Near the potential of zero photocurrent the dependence is close to I* and near the maxima or minima it is close to 1%. Typical results are shown in fig. 5. If we write for the generation of BH. a flux 2fhv which is proportional to I, for a region close to the electrode, where b = [BH-1. Dd2b/aZ2 f 2fbv - 2kRb2 = 0, FIG. 5 pH = mum, -1.01 1 1 I -1.5 -1.0 -0.5 0.0 Wmnad .-Variation of photocurrent with light intensity for 9,lO-anthraquinone 2,6-disuIphonate at 0.80. Each line is displaced on the ordinate by 0.2 log units. Anodic branch, before maxi- 0, gradient = 0.64 ; at maximum, A, gradient = 0.72.Cathodic branch before maximum, 8, gradient 0.66; at maximum 0, gradient = 0.76. The boundary conditions for this equation are and where k' describes the electrochemical rate constant for the destruction of BH- on the electrode. Writing and then b = ubm = u J'vikR, z = xZR = x J D / J k , b , = xJp/(fhDkR)*, (4) (5) d"u/dx"+ 1 -u'" = u, where as x+m, u-1;W. J . ALBERY, M. D. ARCHER, N. J . FIELD AND A . D . TURNER 33 and at x = 0, where 2, describes the reaction layer over which the electrode perturbs the concentration of BH* from its boundary value of b,. Integration of eqn (6) gives From eqn (7) and (9) for A ~ 0 . 1 uo N_ 1 and Whereas for A> 10, uo N 0, and from eqn (3), (7) and (9) The power in I* arises simply from the boundary concentration of BH., which is destroyed by a second-order reaction.The extra power o f t found in the 1 3 arises also from the second-order kinetics but this time in its effect on the characteristic reaction layer thickness 2, through which diffusion to the electrode takes place (see eqn (5)). The results in fig. 5 fit this pattern of behaviour. As 2, from its dependence on k', becomes larger the exponent n in Zn increases from n near 3 to n near 2. Fifthly, by using different filters one can excite CH2 rather than A. When this is done the same types of curve are observed and, in particular, the potential of zero photo-current is the same. (aulax), = ~ L U ~ , (7) (8) A = k'ZR/D = k'/J2D(f,,kR)%. aupx = J(2p)(i - u) 4 2 + u). (9) jB+k'bm = k' J(&/kR)d% (aulax), N 21 J3 = 1.15, j , -+ 1.15D(b,/zR) = 1.15 JZEki(fh,)a oc I*.(10) In this case BH- is formed by CH,-+CH; hv CHZ +A-+2BH*. I I I 1 I I I 2 3 4 5 6 t l m FIG. 6.-Results of laser e.s.r. experiment, showing second-order decay of S, the signal measured in cm of chart paper. Inset shows e.s.r. spectrum, using a modulating frequency of 100 kHz. 56-B34 PHOTOCHEMICAL GENERATION Sixthly, we have investigated our anthraquinone system using a laser e.s.r. system constructed in * The light from a N2 gas laser firing at 100 flashes a second is absorbed by A. After each flash the decay of the e.s.r. signal at a particular field is stored in a fast data store ; to improve the signal-to-noise ratio the results of 4 x lo4 flashes are averaged.The signal decays with second-order kinetics as shown in fig. 6. Using the apparatus in a different mode the e.s.r. spectrum can be obtained. In the presence of both A and CH, (but not A by itself), we observe a broad signal with no hyperfine structure as shown in fig. 6. After many experiments on the same solution of 9, 10-anthraquinone, 1,5-disulphonate only small quantities of hydroxylated anthraquinones could be found. Furthermore, the concentrations of the origirial A and CH, are not significantly depleted. This behaviour is to be expected if the BH. radicals formed by the laser flash are merely disproportionating back to A and CH2. Seventhly, the formation of BH. in this way agrees with other photochemical experiments 4* 9* lo involving hydrogen abstraction by quinone triplet species.The shape of the current voltage curves can be described by solving the transport and kinetic equations for the following model. SOLUTION Tr(a)-fhv+kRb2 = 0, Tr(b) + 2fhv - 2kRb2 = 0, Tr(c) -by kRb2 = 0. ELECTRODE dark light 3 Id 2e+2H++AZCH2, e + H+ + BH-CH2, 4- k' k ; k-; BH.+A+H++e. In eqn (I 1)-( 13), Tr is an operator describing the transport to a rotating disc electrode," a = [A] and c = [CH,]; a2 a Tr = D- a22 +Cz2--; aZ t 3 k' and k' may describe a single step or a pre-equilibrium followed by a rate-determining electron transfer ; fhy describes the photochemical generation and has the form where Zis the intensity of light, q describes the efficiency of the capture of A'$ by CH2 and the exponential term arises from the Beer-Lambert law.For the dark current i,, with Z = 0, we obtain the usual expression for a disc elect rodeW. J . ALBERY, M . D . ARCHER, N. J . FIELD A N D A . D. TURNER 35 where Z , is the diffusion length and is given by,ll W is the rotation speed in Hz and v is the kinematic viscosity. Under these conditions we can neglect the convective term in Tr and we assume, in the simplified treatment presented here, thatfhv is constant. Eqn (12) then beomes the same as eqn (1) with k' = k i + kL in eqn (8) which defines A. 2, = 0.643 W-*v*D*. For the photocurrent we have to solve eqn (12) for small values of z. We then obtain where, from eqn (3), (4), (5) and (8), and uo is given by or Substitution in eqn (19) gives LuO = 1.15(1 - ~ o ) J l +$UO, uo 2: l/(l+i/1.15).where In eqn (21), the dfhY/kR term describes b, the [BH.] produced by the photochemical generation and kinetic recombination. The square bracket in the denominator describes the decay of the photocurrent whenever the current in the dark is a significant fraction of the limiting current. The remaining two terms describe the electrochemistry of the photochemically-generated intermed- iate. On the other hand at some intermediate potential, k i = kL This is the poten- tial where the BH. is equally likely to be reduced or oxidized : It is the same factor as in eqn (18). At extremes of potential 13.9 1 and the two terms reduce to 1.15 D/ZR. and ip = 0. H+ +e + At-BH*+CH2. H + , e The potential of zero photo-current is a direct measurement of the potential where the free energies of the transition states for the first and second electron transfers in the scheme of squares are equal. Its variation with pH i s connected with the route through the scheme of squares. It is the shift with pH of the potential of zero photo-current which is responsible for the chang- ing appearance of the photo-current curves with pH in fig.4. When the zero is symmetrically placed between the half-wave potentials for the reduction of A and oxidation of CH2 an anodic and cathodic branch of the photocurrent is observed. At higher pH the zero shifts under the A-+CH, wave, the cathodic branch disappears and a limiting current for the anodic branch is observed. In fig. 7 we have also plotted, first the variation with pH of E" for the system, i.e., the potential where there Results for the 9,lO-anthraquinone 2,6-disulphonate are shown in fig.7.PHOTOCHEMICAL GENERATION PH Fic. 7.-Plots for 9J0-anthraquinone 2,6-disulphonate of variation with pH of the potential of zero photocurrent E,, A, E” 0 , and E+ for CH2+A, 0. All E are plotted with respect to the saturated calomel electrode. ECEC + t FIG. 8.-Block diagram for 9,lO-anthraquinone 2,6-disulphonate showing variation of k’ and k ’ with pH and potential ; E3, $-wave potential for CH2+A, E* potential of zero photocurrent, En rate-determining transition states.W. J . ALBERY, M . D. ARCHER, N. J . FIELD AND A . D . TURNER 37 would be no dark current for equal concentrations of A and CH2. Secondly, we have plotted the E+ for the oxidation of CH2 ; it is difficult to measure E+ for the reduction of A because H2 is evolved on the Pt electrode.Analysis of the current voltage curves gives the values of a, the transfer coefficients, given in table 1. TABLE VA VALUES OF a RT d l n Iii F dE a = --- A+CH2 0.54k0.02, CH2+A - 0.41 0.01 dark BHn+CH2 0.53+0.01, light BH.+A - 0.42+ 0.01, The values of la1 are all close to 3 suggesting that the rate-determining steps are indeed single electron transfers. For values of la1 close to 3, we can write at the potential E:, of zero photo current where n, and n2 describe the order of the BH. reactions with respect to H+. Then, dE*/d(pH) 21 -(n2-n1)2.303 RT/P. From fig. 7, dE"/d(pH) = 2.303 RT/F = $dE,/d(pH), and so n2 -n, = 2. The value of 2 means that the mechanism must be ECCE involving the El and E6 transition states which differ by 2 protons. This agrees with the half-wave potential for the oxidation of CH2 being independent of pH below 4.Above a pH of 4 there is a change of mechanism to ECEC with E6 being replaced by E4; i.e., the pK of E6 = 3.5. In fig. 7 we have labelled the different regions with the rate determining transition state for the A-CH2 system. At higher values of pH, the A-CH, system becomes reversible and it is then impossible to measure a photo-current since at all potentials the electrode fails to discriminate the photo-chemically generated species. Photo-currents can only be observed when the system is sufficiently irreversible. Fig. 8 is a block which shows the variation of the rate constants k' and k' for the reduction of A and the oxidation of CH2 as a function of potential and pH.The potential E* of zero photocurrent is the hinge which divides the steep planes, where the second electron transfer is rate determining, from the shallower planes, where the first electron transfer is rate determining. Results for 9,lO-anthraquinone 1,5-disulphonate are shown in fig. 9. Here the ECCE mechanism is found at more basic pH and at a pH of 3.9 the mechanism shifts to CECE going through E3 and E6, i.e., the pKA of E3 is 3.9. Returning to eqn (21), we can test whether the [ 1 + (Z,/D)(k' + k')] term does describe quantitatively the decay of the photo-current as the dark current increases. The limiting photo-current is given by eqn (10) and (21) with t -+ + c I : ip,L/PA = jp,L = 1.15 Dh,/Z,.(24)38 r 1.0 - El/2 - c -0- .. -9- - 0 9- - - a V0-E rl, 9 0.5- /o-o- d l 0 .,a 0 /O 4 0 * A M J I I t PHOTOCHEMICAL GENERATION 150 D -too 6 \ 2 -50 .f ' ECCE I CECE I I i/ -0.20 Q2 -0. I k/.. E6 01 I I I I I J 1.0 2.0 3.0 4.0 5.0 6.0 PH FIG. 9.-Plot for 9,lO-anthraquinone 1,5-disulphonate showing variation of potential of zero photo- current with pH. From eqn (18) and (21), Fig. 10 shows the effect of multiplying the observed ip by the correction factor on the right-hand side of eqn (25) to obtain ip,L. A reasonably constant value is found. - * c (1 + Z D ) / D ( k ' + k') = iP,Jip = iD,L/(I'D,L- i D ) . (25) To measure the electrode kinetics of the intermediate BH- species we make Tafel At the potential of plots of the photocurrent. Typical results are shown in fig.11. zero photo-current we obtain an '' exchange current " given by ip,JFA = kkb,, where, at this potential, k; = j ? l = &.W. J . ALBERY, M . D . ARCHER, N . J . FIELD AND A . D. TURNER 39 From eqn (24) and (26), In this expression kR is the only unknown but fortunately it is raised to the power of only 1 /3. We can estimate a value for kR from the laser, e.s.r. experiments. For this we have to estimate the concentration of radicals produced by each laser pulse in order to interpret the e.s.r. signal as a concentration. We have also measured the -7 I I I 3.15 +O.l t0.05 -0.05 -0.1 -0.15 FIG. 1 1 .-Typical Tafel plot of photocurrent at pH of 1.12 ; ip in A and E measured with respect to the saturated calomel electrode. transient behaviour at the semi-transparent electrode.Fig. 12 shows some typical current transients on switching the light on and off. The “ on ” and “ off” transients have different shapes which is to be expected for a system with second- as opposed to first-order kinetics. It takes longer in the dark to remove the last remnants of the radicals by a second-order reaction, than to build them up when the light is switched t/ms curve eventually reaches zero. FIG. 12.-Typical transient curves for 9,lO-anthraquinone 2,6-disulphonate at pH 1 . I . The decay40 PHOTOCHEMICAL GENERATION on. Analysis of the half-times together with the e.s.r. experiments gives an estimate for kR Of kR-7 x lo' M-' S-l. This value is close to the diffusion-controlled limit. obtain from the data and eqn (27) the values of kb given in table 2.Using this value of kR we then TABLE 2.-vALUES OF k& AT THE POTENTIAL OF ZERO PHOTOCURRENT FOR 9,1O-ANTHRA- QUINONE-2,6-DISLJLPHONATE PH k&/,um s- * 0.67 15 0.80 16 0.85 14 1.12 18 We expect k& to be approximately constant since from eqn (22) and (23), These electrochemical rate constants are not as large as one might have expected but since the overall mechanism is ECCE they are describing a CE process in which there is a pre-equilibrium proton transfer followed by the electron transfer. Finally, in general, faster electrochemical rate constants can be measured for photochemically generated species since they are generated close to the electrode and do not have to cross the diffusion layer from the bulk of the solution. Normally, on a rotating disc electrode one can only measure k' when k'- D/ZD ; for the photochemical species the equivalent condition is k'- D/ZR, and when 2, <ZD, larger values of k' can be measured. We thank Dr. McLauchlan for his help with the laser e.s.r. experiments, Dr. E. J. Bowen, F.R.S. for his advice, the S.R.C. for a studentship for N. J. F. and St. Hilda's College for a research fellowship for M. A. Nigel Field died in 1972 and we dedicate this paper to his memory. R. Gill and H. I. Stonehill, J. Chem. SOC., 1952, 1845. F. C. Anson, J. Electrochem. SOC., 1968, 115, 1155. C. F. Wells, Trum. Faraday Soc., 1961, 57, 1703. A. D. Broadbent, Chem. Comm., 1967, 382. J . Jacq, Electrochim. Acta, 1967, 12, 1345. W. J. Albery and M. L. Hitchman, Ring-Disc Electrodes (Clarendon Press, 1971), chap. 5. P. W. Atkins, J. M. Frimston, P. G. Frith, R. G. Gurd and K. A. McLauchlan, J.C.S. Fnraday II, 1973, 69, 1542. F. Wilkinson, J. Phys. Chenz., 1962, 66, 2569. V. G. Levich, Acta. Plzys. Chem., U. R.S.S., 1942, 17, 257. ' P. W. Atkins, K. A. McLauchlan and A. F. Simpson, J. Phys. E, 1970, 3, 547. lo N. K. Bridge and G. Porter, Proc. Roy. SOC. A, 1958,244,259.

ISSN:0301-7249    

DOI:10.1039/DC9735600028

出版商:RSC    

年代:1973

数据来源: RSC

摘要:

Charge Transfer Reactions involving Intermediates formed by Homogeneous Capture of Laser-produced Photoelectrons BY G. C. BARKER,* D. MCKEOWN, M. J. WILLIAMS Atomic Energy Research Establishment, Harwell, Didcot, Berkshire, England A N D G. BOTTURA AND V. CONCIALINI Istituto Chimico c c G. Ciamician ", Universita di Bologna, Bologna, Italy Received 4th June, 1973 Light pulses of short duration produced by doubling the frequency of the output of a Q-switched laser are used to produce electron emission from a mercury electrode. The study of coulostatic photopotentials connected with the scavenging of hydrated photoelectrons in the solution and follow- ing processes can provide information about the chemical and electrochemical reactivity of short- lived intermediates. Results for solutions containing NzO, NO;, NO, or H30 show few abnor- malities connected with the high light intensity.The reduction of adsorbed H atoms on mercury in contact with an acidified solution appears to be a fairly slow process at potentials in the neighbour- hood of -0.9 V against NCE with a transfer coefficient appreciably larger than zero. Unique information about charge transfer reactions involving unstable intermed- iates of electrochemical interest can emerge from studies of photo-currents connected with the homogeneous capture of hydrated electrons (eJ formed by the hydration of photoelectrons near an e1ectr0de.l~~ It is most usual to use unmodulated light, or light modulated at a low frequency or by a long duration waveform ; measurements then tend to refer to a virtually steady state in the thin e& capture zone bounded by the illuminated electrode.This type of experimental approach provides data of value in the context of mechanism identification but there sometimes is a need also for non- steady state experiments which can to some extent resolve the separate steps in a multi-step photocurrent mechanism. Measurements of the coulostatic photo- potential (AE,) produced by light from a short duration (1.5 ps) high pressure xenon spark have given a value for the diffusion coefficient of the OH radical and half-lives of intermediates.2 Korshunov et al. have employed very short duration light pulses from a nitrogen laser for much the same purpose ; in the present paper we consider the comparable reduction in time scale made feasible by the use of second harmonic light pulses derived from the output of a Q-switched laser.EXPERIMENTAL The apparatus, which resembles that used to study the influence of the laser fundamental on a mercury electrode,6 is shown schematically in fig. 1 . The slightly divergent (ca. 0.5") pulsed light beam (ca. 20 ns, < 150 MW cm-2) from a Q-switched (passive switch) ruby laser (Laser Associates Ltd., Rugby, Model 21 1 A) passes through a frequency doubler (Laser 4142 LASER-PRODUCED PHOTOELECTRONS Associates Model 45). The output from the doubler contains a light pulse (ca. 15 ns, f 2 MW cm-2) of wavelength 347 nm and this output passes first through a CuS04 solution filter to suppress completely the fundamental before entering a large screened room which contains the cell and measuring equipment.If the fundamental is to be used, the doubler is removed and a more dilute CuS04 filter is employed. FIG. 1.- solution -Apparatus (schematic) ; L-Q-switched ruby laser, D-frequency doubler, F-CuSO, filter, SR-screened room, D-diaphram, PD-photodiode, SO-storage oscilloscope, C-cell, A-wideband amplifier, 0-oscilloscope, P-polarising circuit. Inside the screened room the light beam is collimated by a diaphram with a circular aperture (7.5 mm diameter) and then is largely transmitted by an inclined glass slide that is part of a " dose " monitor. The integrated signal from a fast photodiode in this monitor is presented on a storage oscilloscope. The transmitted beam is deflected downwards by a 90" quartz prism and travels to the surface of the test electrode via a vertical (9.5 mm diam.) quartz light guide that is part of the cell.Optical alignment can be checked in various ways and the energy distribution at the top of the guide can be monitored with the aid of heat sensitive paper (fundamental) or a phosphor film (second harmonic). The test electrode is a small mercury pool electrode (ca. 10.5mm diam. at the time of the light flash) supplied continuously during an experiment with mercury from a reservoir to avoid surface contamin- ation. Mercury spills over at intervals of 30-45 s into a large annular mercury pool which serves as the counter-electrode and light pulses are produced at a time 10 s after spillover. The cell and its connections have cylindrical symmetry.A gas lock between the cell and the main part of the mercury column isolates this column and minimises the pickup of any high frequency interference radiated by the oscilloscopes when their time bases are triggered. A simple polarizing circuit provides adequate control of the initial cell voltage while per- mitting measurements to be made under virtually coulostatic conditions for small values of elapsed time. Changes in cell voltage produced by the light pulses are recorded with an oscilloscope (Hewlett-Packard Type 180A, amplifier 1803A), the oscilloscope time base being triggered by a photodiode receiving attenuated light reflected from a further inclined glass slide in the main light path. A wideband (200 MHz) amplifier is inserted between cell and oscillope to improve the noise factor (ca.8 dB for frequencies up 35 MHz). This amplifier tends to attenuate frequencies below 100 kHz and its output cable is terminated by a special circuit to eliminate distortion due to this attenuation for values of t, the elapsed time since the mid-pint of the light pulse, < 1.8 ps. The oscilloscope noise is equivalent to ca. 10 pV r.m.s. at the cell terminals and the overall rise time for the electronic measuring equipment (10-90 %) is e l 2 ns. In the work considered later, the peak value of the interfacial temperature excursion for the illuminated part (ca. 50 %) of the test electrode unless otherwise indicated was smallerB A R K E R , M C K E O W N , W I L L I A M S , BOTTURA, C O N C I A L I N I 43 than 10°C, this excursion being estimated from the peak value of the coulostatic potential waveform connected with the heating of the double layer ' 9 2 $ making allowance for attenuation of the signal appearing at the irradiated part of the interface by the unirradiated part of the interface. The cell was cleaned with chromic acid and good quality water and it was invariably steamed just before use.Water was obtained by distillation in an atmosphere of argon from alkaline permanganate solution. AnalaR chemicals, argon and N20/argon mixtures (British Oxygen Ltd.) were used, usually without purification. The cleanliness of the test electrode surface was examined occasionally by studying (measuring potential changes produced by alternating charge injection from a high frequency square wave source) the variation with potential of the differential capacity of the test electrode with and without mercury flowing.No signs of contamination were ever detected with mercury flowing and, even with no flow, the rate of contamination was surprisingly small at all potentials. All experiments were carried out at the temperature of the screened room (20-24°C). Potentials were measured with respect to the counter-electrode which acted as a n.c.e. CORRECTIONS In addition to the potential waveform of interest (AE,) caused by the scavenging of eG and subsequent processes, there is an interfering waveform (AEh) caused mainly by temporal changes in double layer structure resulting from the deposition of light energy at the interface. Coherent Light s o ' t HQ FIG.2.-Current flow induced by illumination of only part of the surface of the test electrode (schematic). For small values o f t the two waveforms are often comparable in size and it is of the utmost importance that the composite waveform (A&+ AEs) observed using the second harmonic light pulse be corrected for the effect of double layer heating.44 LA SER-PROD U CED PI1 OTOE LECTRON S A difficulty in any experiment in which the entire surface of the test electrode is not illuminated is that inevitably the voltage change at the cell terminals tends to be smaller than the change in interfacial potential for the illuminated part of the electrode. Also, due to the flow of current through the solution after the light pulse between solution near the illuminated and dark parts of the electrode (see fig.2), there is a gradual redistribution of double layer charge which citn produce appreciable distortion at large t of the shape of the voltage waveform at the cell terminals. Usually it is desirable that the maximum value of t should be small compared with the " time constant " for the complex process of charge re- distribution. EVALUATION OF AEh AEh corrections were obtained by essentially the method used in earlier work.2 Studies were made for each solution of the variation with time of A& for constant light dose and various potentials, using the fundamental pulse. For each potential, data from at least six oscilloscope traces were averaged. Similar studies were made of the composite signal, AEh+AE,, using second harmonic pulses.The dose monitor sensitivity varied with wave- length but the relative sensitivities for the two wavelengths could be found by comparing the signals at a potential sufficiently positive (- 0.3 or - 0.4 V against NCE) to make A& (second harmonic) trivial in comparison with A&. To minimise any error connected with slight non-linearity in the double layer, as regards the variation of potential with temperature at constant charge density, as far as possible (the short term variation in dose was ca. +25 %) the energy deposited at the two wavelengths was made equal, CORRECTION FOR CHARGE REDISTRIBUTION We give later data only for t < 0.4 ps and for such values oft any distortion of the observed waveforms was small (the time for attainment of 50 % of the charge redistribution occurring after a potential step appears at the illuminated part of the interface is of the order of 30 ps).Our results for a variety of systems suggest a slight decay in waveform with increasing t corresponding approximately to an exponential decay of time constant 4 p s for the time interval t = 0.1 p s to t = 0.3 ps. Though the application of small corrections based on this somewhat uncertain experimental fact is difficult to justify completely, we have, in figures appearing later, applied tentative corrections for t G0.4 /is that are the inverse of distortion resulting from the passage of the AEs waveform through a condenser-resistoi coupling of time constant 4ps. The corrections although small, if seriously in error, could affect appreciably some of the conclusions drawn from results for solutions containing H30f.While it would be advantageous from several viewpoints to use a DME rather than a pool, due to the reduction in electrode size errors due to charge redistribution would grow if a DME were employed unless some way of producing uniform illumination of the mercury drop can be found (e.g., by multiple reflections within a spherical reflector surrounding the DME). RESULTS AND DISCUSSION We give here only preliminary results for solutions containing N20, NO;, NO, and H30+. The main aim of the experiments has been to show that high light intensities provided by a pulsed laser can give reliable results that are in harmony with what is already known or believed about light-induced electron emission and subsequent processes.The units used for A& are arbitrary but invariant and AEs invariably refers to a constant light " dose ". NITROUS OXIDE SYSTEMS The capture of e& by N,O results almost instantaneously in the deposition of OH These radicals are not of outstanding electrochemical radicals in the solution1*BARKER, MCKEOWN, WILLIAMS, BOTTURA, CONCIALINI 45 interest but they probably are formed in the reduction of H202 at mercury (possibly also at other electrodes) taking the view that the rate-controlling step is then H202+e- -+ H202 (bet) ( a ) H,O; -+ OH-+OH (horn?) (b) this being followed by and OH+e- 3 OH- (het). (4 Reaction (c) is part of the N20 photocurrent mechanism and is known from the influence of aliphatic alcohols on the N20 “ steady state ” photocurrent to be virtually diffusion-controlled.The diffusion coefficient of OH (DOH) has been determined approximately by studying the variation of the coulostatic photopotential with time at large values of t using light from a xenon spark source.2 Though special waveform shaping circuits were then inserted in the photopotential measuring circuit, to elimin- ate the adverse influence of the long duration tail of the light pulse provided by the spark, it was not possible to detect easily an initial fall in photopotential magnitude shortly after the light pulse when the N20 concentration was small that was connected with the rapid return of unscavenged eG to the electrode. To be absolutely certain about the various steps initiated by light absorption (probably homogeneous absorp- tion) it is desirable to check the variation of photopotential with time with and without scavengers present ; for this and other reasons the N20 system has been reinvestigated.The photopotential observed in the absence of a scavenger resembles that observed earlier with the spark source,2 provided that the “dose” is low enough to avoid significant reaction of e& in pairs, a requirement impossible to reconcile with a good signal : noise ratio. Results for a low dose (with it will be noted the electric vector virtually parallel to the electrode surface) point to equality for the magnitudes of AEh and AE,(C, = 0) at a potential in the vicinity of - 1.5 V against NCE, a result in agreement with earlier work.2 Results for somewhat larger doses, but corresponding to a peak temperature excursion + 20°C at the irradiated part of the interface, suggest a rate constant in 1 M KCl for in the vicinity of 2 x 1O1O M-I s-l a finding in harmony with the literature after taking account of the change in ionic strength.Photopotential curves for 1 M KC1 saturated with 10 % N20/90 % argon are given in fig. 3. These show the expected initial decline in AEs shortly after the light pulse which precedes the gradual increase in AE, with increasing t connected with relatively slow diffusion and eventual reduction of OH radicals. The theoretical curve for a light pulse of 15 ns duration obtained using an integrated form of the theoretical expression for the coulostatic photopotential 2 p (the expression for the photocurrent at constant E produced by a light step) ignoring distributed deposition of e,g.and assuming (c) to be diffusion-controlled does not agree perfectly with the experimeiital curve for - 1.3 V at small values of t . The theoretical curve in fig. 3 refers to a value for the ratio DoIl/(DeG- DOM) of 0.5 and the value 1.5 x lo7 s-I for the product of electron scavenging rate constant k, and scavenger concentration C,. In practice the minimum in the curve occurs at a value oft consistent with k,C,r0.9 x 1 O7 s-’ rather than the assumed value. We incline towards experimental error (other than that caused by the finite rise time of the electronic apparatus) as the cause of the46 LASER-PRODUCED PHOTOELECTRONS difference in curve shape, though the shape could have been slightly influenced by a finite life for N20-, by distributed deposition of e; or by some kinetic control of reaction (c).It should be mentioned that the value of k, obtained by comparing (a) the absolute magnitude of AEs(Cs = 0) for known t and a low dose with (b) the corresponding value of AEs for t+co obtained (by extrapolation using theory) for 1 M KCl saturated with 10 % N20, taking account of the fact that an appreciable fraction of the electrons are captured, is 5.0+ 1.4 x lo9 M-l s-l in agreement with the literature value lo (low ionic strength) of 5.6 x lo9 M-' s-' . The apparent threshold potential ~5':~ for electron emission suggested by a plot of A*, the charge lost ( t = 0.3 ps), in the coordinates - E is ca.0.1 V against NCE, in fair agreement with the value expected for 347 nm from the work of Pleskov and co-w~rkers.~ -1.3V \- -l.IV -0.9 V -0.7V Light Pulse 0 0. I 0 -2 0.3 0.4 t ime/ps FIG. 3.-Variation of AE, with time for 1 M KCl saturated with 10 % N20/90 % argon ; dashed curve is corrected for charge redistribution ; x-theoretical variation for 15 ns light pulse, DOH/ DOH) = 0.5 and k,C, = 1.5 x lo7 s-' (uncorrected for finite rise time of electronic equipment). Using the laser light source it has not proved possible so far to determine DoFr with improved precision. This mainly reflects the relatively poor reproducibility of the laser pulse as regards both (short term) intensity and (long term) energy distribution which overrides the benefit gained from the well-defined pulse shape.OXIDATION OF co, The introduction of formate ion in the N 2 0 system throughout much of the acces- sible potential range produces partial suppression of the " steady state " photo- current 1 1 9 l 2 due to the homogeneous consumption of some OH by OH+HCO; -+ COY +H,O (horn) (4 which tends to be followed by kf COY .+ C02+e- (het). Reaction (f) we have found too rapid to study using light pulses of duration 100 {isBARKER, MCKEOWN, WILLIAMS, BOTTURA, CONCIALINI 47 or longer; to obtain more information about its kinetics we have examined AEs using 1 M KC1 that was saturated with 10 % N 2 0 and contained also HCO,. Some typical corrected results are given in fig. 4. The slow rise in A& found for N 2 0 when t > 0.1 p s now tends to be replaced by a decay in A& CO; oxidation is quite 3 0 I .- ; I \ - - -1.5V I 0 n 0 - L45V r\ 0 I I 1 0 0.1 0.2 0 .3 0.4 timelps FIG. 4.-Variation of AE, with time for 1 M KCl+5 x 90 % argon. M HCO, saturated with 10 % NrO/ slow in the potential range - 1.5 to - 1.3 V where normally, in a " steady state " experiment, the photocurrent would be reduced by a factor close to 10". The results in fig. 4 point to a value for kf at - 1.3 V against NCE of ca. 0.4 cm s-'. OXIDATION OF THE a ETHANOL RADICAL EtOH, like HCO;, produces over much of the potential range partial suppression of the " steady state " N 2 0 photocurrent. In this case, some of the OH radicals are converted homogeneously to a ethanol radicals which, after diffusion to, and adsorp- tion on the electrode, are oxidized to acetaldehyde.The time constants for their oxidation (0.2 M KCl pH ca. 8) are 1.7, 1.25, 0.65 and ca. 0.3 ms respectively at potentials of - 1.6, - 1.5, - 1.4 and - 1.3 V against 0.2 M CE. Recent pulse radio- lysis work l4 has suggested a fairly rapid change in oxidation rate constant with potential at more positive potentials, pointing perhaps to a change in mechanism. However, the results in fig. 5 show that for " neutral " 1 M KC1, even at - 1.0 V against NCE, the oxidation of the alcohol radical is so slow a process that the values of A& ( t = 0.3 ,us) for e 4 capture alone calculated from data obtained in the absence of EtOH, making allowance for incomplete homogeneous capture of OH by EtOH, differ from the experimental values for the EtOH-N20 system by no more than the experimental error.* No marked acceleration in alcohol radical oxidation was * The scatter in the points is unusually large in this case.48 LASER-PRODUCED PHOTOELECTRONS I 5 v) Y .d d 3 . i ' O - ry" a 5 - observed evenat - 0.68 V against NCE* and it thus seems likely that the rate controlling step is the chemical step - H20 + CH,CHOH+CH3CHO-+ H30+ I I Hg Hg this being followed by the rapid transfer of an electron to the electrode and the desorption of CH3CH0. 2 0 r - I.OV LL 0 0. I 0.2 0.3 0.4 time/ ys FIG. 5.-Variation of AE, with time for 1 M KCI+0.2 M EtOH saturated with 10 % N20/90 .% argon ; 0, values at t = 0.3 ps calculated from data for 1 M KCl saturated with 10 % N20/90 %argon assuming no oxidation of the ethanol radical and 86 % homogeneous capture of OH by EtOH.NITRATE AND NITRITE SYSTEMS The behaviour of these systems in " steady state " photocurrent experiments is now fairly well understood 2* and, in the present work at low light dose, the behav- iour of NO; ( 5 x M) in 1 M KC1 was much as expected. Due to the quite slow reaction of NO $ - with water,' AEs was governed almost entirely by electron capture ; experimental data plotted in the coordinates Aq;/'E gave almost linear plots and E& values close to 0.0 V against NCE. Results for NO; in the same medium pointed unambiguously to a half-life for the electron adduct (NO;- ?) of the order of 15 ns as the initial decay in AE, arising from the return to the electrode of uncaptured eG was not observedwith [NOT] = 5 x 10-3M but there was an abnormally sudden increase in AEs in the vicinity of - 1.1 V for all values of t between 0.05 and 0.3 ps.Results for this system indicate that the half-life at the electrode surface of NO formed by reaction of the electron adduct with water must become very short in the same region of potential. M and 5 x REDUCTION OF ADSORBED HYDROGEN ATOMS " Steady state " photocurrent measurements using acidified aqueous solutions l * and acidified aqueous-ethanolic solutions 1 * l 3 have shown that H atoms formed * At this potential AEh vanishes if the dose is small.BARKER, MCKEOWN, WILLIAMS, BOTTURA, CONCIALINI 49 homogeneously by capture of e i by H30+ are strongly adsorbed on mercury and, when adsorbed, tend to vanish rapidly by oxidation to H30+ at positive potentials and by reduction to molecular hydrogen at negative potentials. The latter step, according to the work of Pleskov and co-w~rkers,~ involves a water molecule at high pH or a hydrogen ion if the pH is low (< ca.3). A preliminary attempt,, using light from a zenon spark, to determine the time constant for the reduction of adsorbed H (1 M KC1+ M HCl) gave an approxi- mate value of 2 p s at - 1.0 V against NCE. More recent results using the laser light source are in approximate agreement with this result. Fig. 6 and 7 show for different values of E the variation of AE, with time for 1 M KCl containing HCl at the concentrations 0.01 M and 0.1 M respectively. Of particular importance from an interpretative viewpoint are the relatively large increases in AE, for t = 0.3 ,us on changing E from - 1.1 to - 1.3 V in the case of the less acid - 0 s v 0 a -0.7 V I I I 0 .2 0.3 0.4 timelps FIG. 6.-AEs for 1 M KCI+O.Ol M HCl ; dashed curves are corrected for charge redistribution ; 0, calculated values if H atom reduction were diffusion-controlled. solution, and a similar anomalous change in the case of the more acid solution on changing E from -0.9 to - 1.1 V. The potential at which half the adsorbed H atoms are reduced is in both cases somewhat positive with respect to the potential region associated with the large increase and, ruling out the possibility of experimental error (experiments were repeated several times with unusually repro- ducible results), it would seem that these large increments in AEs are caused by the effect of potential on the rate constant of (h) e- + H(ads) + H30+-+H2 + H,O.50 LASER-PROD U CED PHOTOELECTRONS 0- v - 1.3v P - I .l V O I A n n v v - 0 . 7 V I I I 0.1 0.2 0 . 3 0.4 time/p FIG. 7.-Data for 1 M KCl+O.l M HCl corresponding to that in fig. 6. The potential variation for t > 0.1 ps and E = - 1.3 V is for both solutions close to that expected for diffusion-controlled reduction of H atoms (assuming DH = 1.5 cm2 s-l and also ks = 6 lo9 M-l s-l f or ea; capture by H30+ in 1 M KC1).l6 Making use of this fact, the Brodsky-Gurevich Law l7 and the apparent threshold potential for 347 nm calculated from the data of Pleskov and co-worker~,~ the value of AE, corresponding to diffusion controlled consumption of H for - 0.9 and - 1.1 V, and with t = 0.3 ps, can be calculated.Comparisons of the observed gradual rise of AE, with these calculated values leads to time constants for the reduction of adsorbed H, some values of which are given in table 1. The tentative values for TABLE TI TIME CONSTANT FOR THE REDUCTION OF ADSORBED H ATOMS time constant potentiall V against NCE 1 M KCI+O.l M HCI 1 M KCl + 0.0 1 M HCI - 1.3 - 1.1 - 0.9 <20 ns -30 ns 0.5 p s -70 ns 0.4 ps >l ps - 1.3 V have been estimated from the shape of the (AEs, t ) curves for t <O. 1 ps. The reliability of all these values is difficult to assess but these results agree qualitatively with the earlier value for - 1.0 V and 1 M KCI, M HCI. Also they show approximately the correct dependence on [H,O+]. The decrease in time constant between - 0.9 and - 1.1 V is partly attributable to the increase in surface concentra-BARKER, MCKEOWN, WILLIAMS, BOTTURA, CONCIALINI 51 tion of H30+ but it seems that the apparent transfer coefficient of reaction (h) may be appreciably larger than zero.The present work suggests that a Q-switched laser is a convenient source of inter- ference-free light pulses of short duration and " clean " shape. Elsewhere we shall show that the present experimental method is of value in clarifying complex effects observed when e; is captured by the divalent ions of Zn, Mn, Ni, Co and Fe.I8 We are grateful to Messrs Luke Randolph and Peter Taylor for experimental help. G. C. Barker, A. W. Gardner and D. C. Sammon, J. Electrochem. Soc., 1966, 113, 1182. G. C. Barker, Ber. Bunsenges. phys. Chem., 1971, 75,728. Yu. V. Pleskov and Z. A. Rotenberg, J. Electroanal. Chem., 1969, 20, 1. Z. A. Rotenberg, V. I. Lakomov and Yu. V. Pleskov, J. Electroanal. Chem., 1970, 27, 403. L. I. Korshunov, Ya. M. Zolotovitskii and V. A. Benderskii, Uspekhi Khim., 1971,40, 1511. G. C. Barker, G. Bottura and A. W. Gardner, J. Electroanal. Chem., in press. M. S. Matheson and J. Rabani, J. Phys. Chem., 1965, 69, 1324. G. C. Barker and A. W. Gardner, J. Electroanal. Chem., in press. l o M. Anbar and P. Neta, Int. J. Appl. Rad. Isotopes, 1967, 18, 493. '' G. C. Barker and G. Bottura, J. Electroanal. Chem., in press. l 3 G. C. Barker, Electrochim. Acta, 1968, 13, 1221. l 4 M. Gratzel, K. M. Bansal and A. Henglein, Ber. Buizsenges, phys. Chem., 1973, 77, 11. l 5 G. C. Barker, P. Fowles and B. Stringer, Trans. Faraday SOC., 1970, 66, 1509. l 6 G. C. Barker and V. Concialini, J. Electroanal. Chem., in press. l 8 G. C. Barker, V. Concialini and D. McKeown, unpublished work. ' L. I. Korshunov, Ya. M. Zolotovitskii and V. A. Benderskii, Elektrokhim., 1968, 4,499. D. J. Schiffrin, private communication. Yu. Ya. Gurevich, A. M. Brodskii and V. G. Levich, Elektrokhim., 1967, 3, 1302.

ISSN:0301-7249    

DOI:10.1039/DC9735600041

出版商:RSC    

年代:1973

数据来源: RSC

摘要:

Investigation of Intermediates by Electron Photoemission from Metal into Electrolyte Solution BY Yu. V. PLESKOV, Z. A. ROTENBERG, V. V. ELETSKY, AND V. I. LAKOMOV Institute of Electrochemistry, Academy of Sciences of the U.S.S.R., Moscow Received 11 th June, 1973 During electron photoemission from an illuminated electrode into an electrolyte solution con- taining scavengers of hydrated electrons, stable and unstable products from reduction of the scaven- ger are formed in the vicinity of the electrode, which, at the same time, are intermediates of some electrochemical reactions. Thus, electron photoemission into acid solutions produces atomic hydrogen, that into nitrate solutions-the ion-radical NO;-. In this type of experiment photo- emission can be used, on the one hand, as a convenient source of intermediates and, on the other, as an instrument for measuring the rates of their chemical and electrochemical reactions.Investigation of the simultaneous ionization and reduction of atomic hydrogen at a mercury electrode has shown one of the two reactions to be activationless (the transfer coefficient is zero). The rate constant of the decomposition of the unstable ion-radical NO$- in the solution bulk has been measured, as well as the transfer coefficient of its anodic oxidation on mercury. The photoemission of electrons from metal into solution is accompanied by a number of processes associated with formation and further transformation of the solvated Among these, of great interest are the homogeneous and het- erogeneous reactions involving free radicals.The kinetics of these reaction affect the value of the measured photocurrent so that photoemission may be used as a convenient technique of quantitative study of such processes. The experiments described below were conducted, mainly using a mercury electrode which was illuminated by ultra-violet light from a high-pressure mercury lamp. The experimental techniques are described in detail in ref. (2) and (3). KINETICS OF ELECTROCHEMICAL TRANSFORMATION OF ATOMIC HYDROGEN The properties of atomic hydrogen have attracted the attention of scientists for a long time. This is mainly due to the fact that the atomic hydrogen is an intermediate product of one of the best-studied reactions : electrochemical hydrogen evolution. Hydrogen atoms are formed during photoelectron emission into acidic aqueous solutions as a result of a capture by hydrogen ions of hydrated electrons : Hydrogen atoms either recombine or approach the metal surface and enter into electrochemical reactions.The first mechanism of hydrogen removal from solution is, apparently, not effective since recombination, being a second-order reaction, is very slow when the concentration of hydrogen atoms is low. Therefore, practically all the atomic hydrogen is removed by an electrochemical mechanism. show that, in a wide range of potentials, e,; + H,O++H + H20. (A) Studies using a mercury electrode 52YU. V. PLESKOV, Z. A. ROTENBERG, V. V. ELETSKY, V. I. LAKOMOV 53 hydrogen atoms preferentially enter into a reaction of the electrochemical desorption type H + HS + e--+H, + S-.(B) In the above equation, HS is a source of protons, that is, hydrogen ions and water molecules in aqueous solutions. At more positive potentials, as well as this process, there is a possibility of a reaction of atomic hydrogen involving ionization H + H20+H,0++e- (C) which brings about a decrease of the measured photocurrent. This reaction was detected in experimental studies of photoemission from mercury by Barker et al. and quantitatively studied in detail in ref. (4). It should be noted that, in this case, photoemission is an effective source of atomic hydrogen near the electrode surface and that such a source proved to be more convenient than the earlier ones 5-7 because of its simplicity and the possibility of con- trol of hydrogen influx rate which it provided.KINETIC EQUATIONS FOR THE ATOMIC HYDROGEN REMOVAL Taking into account the reactions (B) and (C), we may define the experimentally measured photocurrent j by the following equation : + t j = h + j ~ - j ~ wherej, is the emission current Iless the current I.. produced by return to the electrode of the electrons which were not trapped by scavengers in the solution ; jH is the cathode current of the reaction (B) andjH is the anode current of ionization (C). Essentially, there are two possible mechanisms for the reactions (B) and (C)-one involving adsorption of hydrogen atoms on the surface, and the other directly from solution, by-passing the adsorption stage. If the reaction proceeds simultaneously by both mechanisms the following relations are valid under stationary conditions : -* 4- - t i - + c j = j , + (k, - kI)cH(0) + (k2 - k2)0. (24 In the above relations, kads is the adsorption rate constant expressed, as the other kinetic constants, in electrical units; to a first approximation it does not depend on the potential; k, and k, are, respectively, rate constants for cathode and anode removal of atomic hydrogen directly from the solution; k , and k2 are the same constants for adsorbed hydrogen ; ~ ( 0 ) is concentration of hydrogen atoms near the electrode, and 0 is the surface coverage by adsorbed hydrogen atoms.The rela- tions (2) do not take into account recombination of adsorbed hydrogen atoms on the electrode surface : this is justifiable for a mercury electrode5 If the reaction (B) proceeds with simultaneous participation of water molecules and hydrogen ions, the rate constants k, and k2 may formally be divided into two components: ki = kl+ Xky, where kf and ky are, respectively, rate constants for hydrogen removal + 4- 4- -+ -+ -+ -b + - + + +54 ELECTRON PHOTOEMlSSlON involving water molecules and hydrogen ions and Xis the molar fraction of H30+ in the solution.By eliminating 8 and cH(0) from the relations (2) we obtain the following equation for the measured photocurrent : r + -+ 1 kl kzkads I + + + -+ c + c j = 2j0 (3) In the limiting case, when the oxidation reactions may be neglected, all the atomic hydrogen is reduced on the electrode so that the measured photocurrent is twice the emission current ( j = 2j0) and its relationship with potential 4 is defined by the “ 5/2 law ” : joc($-$o)5/2, where #o is the photoemission threshold.In the general case, since the constants in eqn (3) are themselves functions of the potential, the relationship between j and 4 is complicated and does not allow for direct experi- mental verification. It is expedient to consider the limiting cases for comparison between theory and experiment. (a) The reactions (B) and (C) proceed by-passing the adsorption stage (/cads = 0). Then -+ k1 j = 2j0 . t kI+ kI + t The relationship between constants ki, ki and potential may be represented by the following equations : ki = klo exp( - a,F+/RT), & = kc,, exp(PiF4/RT), + + where ai and Pr are the transfer coefficients for the appropriate reactions. Then the relationship between j and 4 may be written as (b) Atomic hydrogen enters into both reactions kl = 0).From the eqn (3) we obtain 4- (4) -+ only as an adsorbed particle (k, = (c) The reaction (B) proceeds through the adsorption stage while the reaction (C) proceeds directly from the solution (k, = k2 = 0), then - + c + = - 2.3RT “og7+log*]. 2jo-j k1o PIF ( d ) The reaction (B) proceeds bypassing the adsorption stage while the oxidation of hydrogen proceeds via the absorption stage (k, = kI = 0). 4 4 - In this case we obtainYU. V. PLESKOV, Z . A. ROTENBERG, V . V. ELETSKY, V . I. LAKOMOV 55 The eqn (4)-(7) exhaust the real mechanisms for removal of atomic hydrogen. These equations are similar to the kinetic equations for slow discharge, with the only difference that instead of discharge current we have dimensionless parameter (2j0 -j)/’ which is equal to the ratio of oxidation current jH to reduction current j,.The structure of the eqn (4)-(7) provides for linear relationship between 4 and 10g(2j0 - j ) l j with the line slope defining transfer coefficients for the corresponding reactions. -+ c RELATIONSHIP BETWEEN T H E KINETICS OF HYDROGEN REMOVAL, POTENTIAL A N D COMPOSITION OF THE SOLUTION Fig. 1 presents the (photocurrent, potential) relationship for a mercury electrode plotted injoa4, 4 coordinates for three solutions with various acid concentrations, but with constant (1 N) overall electrolyte concentration. For more acidic solutions the potential, which corresponds to the beginning of the departure from linearity arising from hydrogen oxidation, is somewhat shifted to the more positive values.With specific adsorption of halogen ions, particularly iodine, the reaction of hydrogen oxidation is accelerated in relation to the reaction of its cathode reduction. Adsorp- tion of tetrabutylammonium cation is accompanied by a reverse effect: drop of photocurrent in the potential region studied disappears. 2 x ‘I I -0.5 - 1.0 4lV FIG. l.-(joe4, 4) plot for the mercury electrode in K2S04+ H2S04 solutions. Concentrations of H30+ : 1-10-3, 2-10-2, 3-lo-’ N. Current is expressed in arbitrary units, potential is determined with respect to the saturated calomel electrode. On a bismuth electrode, the drop of photocurrent associated with hydrogen oxidation begins at more negative potentials (-0.8 V) which may be due to a higher energy of hydrogen adsorption than in the case of mercury.Photocurrent decay was not found on a lead electrode (at 4 < -0.7 V) and on a cadmium electrode (at For the experimental verification of the eqn (4)-(7) it is necessary to know, besides current j , also current 2j0 in the absence of atomic hydrogen oxidation. This current may be determined by extrapolating the linear portion of the joS4, 4 plot into the potential range where experiments show departure from linearity, or by measuring photocurrents with other scavengers whose products of interaction with hydrated electrons are not oxidized on the electrode. This requirement is met, for example, by 4 < -0.9 V).56 ELECTRON PHOTOEMISSION nitrous oxide which produces after electron capture, the OH radical.This radical is reduced on the electrode in the whole potential range under consideration. Fig. 2 presents the experimental (4, 10g[(2j0 -jib]) plots which are, actually, straight lines for all the solutions studied thus indicating the validity of the eqn (4)-(7). > 2 -0.5 - 1 L I 0 I log(2jo -.M FIG. 2.-( 4, 10g(2j0-j)/j) plot for the mercury electrode in various solutions. The concentration of H30+ is 0.01 N. 1,0.9 N K,SO,; 2,O.l N KCI; 3,O.l N KBr; 4,O.l N KT. THE MECHANISM OF ATOMIC HYDROGEN REMOVAL On the basis of the above information we concluded that the mechanism of hydro- gen removal via the adsorption stage is the most realistic since the reactions of ioniza- tion and reduction of hydrogen directly from the solution in the studied potential range should be activationless (that is, their rates do not depend on the potential) and, thus, cannot explain the photocurrent decay observed in the experiments.A similar conclusion was made by Barker who studied photocurrents in acidic solu- tions in the presence of ethanol. We now consider the influence of pH on the photocurrent. For the quantitative study of the effect of pH it is convenient to introduce the potential 4* at which cathodic current is equal to anodic current, that is, (2j0 - j ) / j = 1. The potential 4* may be determined by the point of intersection of the nonlinear portion of the experi- mental curve ( j o e 4 , 4) with a straight line passing through the extrapolation point (&) with slope which is 20a4 times less than the slope of the experimental curve.Fig. 3 presents $* as a function of concentration of hydrogen atoms c ~ ~ ~ + in the concentra- tion range from 0.001 to 1 N. The potential 4* is practically independent of pH in solutions with low acidity (<0.01 N). When acidity of the solution is increased, the potential 4* rises steadily, and its greatest change is observed for the high values of concentration cH30+. The marked influence of 4* on pH indicates that the reaction of electrochemical desorption simultaneously involves both hydrogen ions and water molecules. Krishtalik lo was the first to suggest the possibility of adsorbed hydrogen removal by electrochemical mechanism via water molecules in acidic solutions. The observed ($*, pH) dependence might formally be explained by suggesting intermediate formation of H i radical in acidic solutions ; this radical is then electro- chemically reduced.However, the reaction rate constant for the interaction between H and H30+ is so low that this reaction’s contribution may be neglected.YU. V. PLESKOV, Z . A. ROTENBERG, V . V. ELETSKY, V . I. LAKOMOV 57 - 0.61 I i * I * I - 0.5 I > 1 ,c2 16’ t CH30+ /mol I.-’ FIG. 3.-Variation of (b* with concentration of H30+ (mercury electrode). Thus, the reactions involving atomic hydrogen are represented by the following scheme : It follows from the slope of the 4, 10g[(2j0-j)/’] straight lines of fig. 2, that is equal to 100-120 mV, that a+p = + for two simultaneous reactions (B’) and (C’) if eqn ( 5 ) is valid.The transfer coefficients for electrochemical reactions involving adsorbed atomic hydrogen are usually equal to 3. Hence, the observed value of 3 for the sum of transfer coefficients for the two reactions indicates that one of these reac- tions is activationless (transfer coefficient is zero), that is, its rate is potential-independ- ent. It has not yet been decided definitely which one of the reactions is activationless. There is some evidence suggesting the activationless character of the first as well as of the second reaction. Thus, it should be noted that the ionization of adsorbed hydrogen is the reverse reaction in relation to the discharge of hydrogen ions. If the ionization is activation- less, the discharge of H,O+ should proceed in the same potential range as a barrierless reaction. According to Kri~htalik,~.O who studied the barrierless discharge of hydrogen on mercury, the transfer from the usual discharge to a barrierless one occurs at the potentials more positive than -0.5 V. Therefore, in the potential range where we observe the photocurrent decay, the ionization should be an activation reaction, so that the electrochemical desorption should be activationless. However, direct comparison of experimental data for hydrogen ion discharge with the results of photoemission experiments is not rigorous since in these two cases the hydrogen at oms are probably not energetically equivalent. The comparison of photocurrent magnitudes at -0.4 V for the two reactions studied, shows that at this potential the electrochemical desorption is tens of times slower than the hydrogen oxidation.At the more positive potentials the photocurrent58 ELECTRON PHOTOEMISSION practically disappears, that is, all the atomic hydrogen is oxidized. This ratio of rates for the two reactions indicates that the activationless reaction is the hydrogen ionization rather than desorption. The specific adsorption of ions does not influence the rate of activationless ionization since the latter with p = 0 does not depend on the structure of the electric double layer. For the activated electrochemical desorption, the change of $'- potential due to adsorption of, for instance, iodine, gives rise to two opposite effects : decrease of the discharge rate and increase in the vicinity of the electrode of hydrogen ion concentration involved in the process.The experimental results demonstrate the prevalence of the former effect. This is due, particularly, to the participation of water molecules in the reaction of electrochemical desorption. Although we were not able to establish which one of the two above reactions is activationless-the ionization of atomic hydrogen or its reduction, the very fact of experimental observation of an activationless process is of great importance by itself. It should be noted that the concept of activationless processes has existed in electro- chemistry for a comparatively long time 11* l 2 but up till now it was not possible to observe directly the limiting currents of activationless discharge. For this observation it is necessary to overcome diffusion limitations which are essential for high discharge rates.The photoemission permits one to by-pass these difficulties, since in this case we measure not the rate of an individual process, but the ratio of rates for two con- current processes involving the same material (atomic hydrogen) so that the diffusion difficulties are mutually cancelled out. ELECT R 00x1 D AT1 ON AND H 0 M 0 GENE 0 US D E CO M PO S I T TO N The chemical reactions accompanying capture by anion NO, of a hydrated OF ION-RADICAL NO$- electron may be represented by a following scheme : k r N O j +e&-+NO:- (in solution) kv NO:-+ H,O-+NO; + OH+OH- (in solution) keA NO $ -+NO, + e- (on electrode) OH + e-+OH- (on electrode). (GI The oxidation of NO:- is accompanied by a photocurrent decrease.'.9 * 13-15 the measured photocurrent with electron capture by NO; : where ZA is the oxidation current for NO;-. The factor 2 indicates that each act of decomposition of NO;- in the solution is accompanied by transfer across the interface of one additional electron due to OH reduction. The oxidation current is determined by the relation IA = k e A c ( 0 ) , where c(0) is the concentration of NO:- in the vicinity of the electrode, keA is the oxidation rate constant for NO:-, which is potential- dependent as for any electrochemical reaction : Diffusion and homogeneous reactions of formation and decomposition of the anion- radical NO : - may be represented by the following equation : Taking into account the reactions (D)-(G), we obtain the following equation for j = 2(I- 1, - I A ) (8) k e A = k t ~ exp(PF$/RT). Dd2c/dx2 - kvc + kACACe = O (9)Y U .V. PLESKOV, Z . A. ROTENBERG, V . V. ELETSKY, V . I. LAKOMOV 59 where D is the diffusion coefficient for NO:-, k , is the rate constant of NO$- homo- geneous decomposition, CA and kA are NO; concentration and its rate of reaction with hydrated electrons of concentration c,. Thus, the product kAcAce defines the rate of NO:- formation in solution. This equation, together with a similar equation for hydrated electrons, allows one to determine the current I A . Integration of eqn (9) with boundary conditions c(00) = 0, keAC(0) = D(dc/dx),,o yields the following expression : where Q = (kAcA/De)) and Q, = (kv/De)* ; the function @(x) defines deposition rate of hydrated electrons in solution, and De is the diffusion coefficient of hydrated electron.The final expression for the photocurrent is obtained ‘ 9 l6 by substituting (10) into (8) and taking into account the expression for I, (for the case of infinitely fast capture of hydrated electrons by the electrode surface) : At very negative potentials when keA < QvD (that is, NO:- oxidation may be practically neglected) j = 2(I-Ie)= In the reverse limiting case (keA>> QvD) we obtain dx} . In both cases the photocurrent is proportional to the emission current, that is, its potential dependence is defined by the “ 5/2 law ”, to a first approximation. The experimental study of this reaction demonstrated both of the discussed limiting cases on the actual (photocurrent, potential) curve for mercury and bismuth elec- trodes 13* l4 (fig.4 and 5). The linear portions at high and at low potentials in 4 FL - - 2 - - I - r I H’ - 0.5 - 1.0 - 1.5 4lV FIG. 4.-(j0-4, 4) plot for the mercury electrode in 1 N KN03 solution (1) and the same solution with addition of tetrabutylammonium bromide (2).60 E LECT R 0 N P H 0 TOE M I S S I 0 N ( j o a 4 , 4) coordinates are separated by the region of monotonous decay of the photo- current (from -0.7 to - 1.2 V) as may be seen on fig. 4. Addition to the solution of small amounts of tetrabutylammonium ( mol l.-I) produces a characteristic change of the shape of the curve. At potentials more positive than - 1.4 V (desorp- tion potential of tetrabutylammonium) the curve becomes practically a straight line.FIG. 5.-(jos4, - 0.5 -1.0 4lV 9) plot for the bismuth electrode in 0.5 N solution of KN03. It is evident that the adsorbed tetrabutylammonium produces two concurrent effects : emission current decrease and inhibition of NO:- oxidation (cf. inhibition of hydrogen oxidation discussed above), thus resulting in absence of photocurrent decay at positive potentials. Experimental data used with expression (12) made it possible to evaluate the rate constant k, for homogeneous decomposition of NO:- and the transfer coefficient p for its clectrooxidation. The rate conslant k , was deter- mined in sufficiently concentrated (0.5-1 .O N) nitrate solutions where practically all the hydrated electrons were captured. If the rate of NO:- decomposition in the solution is not too high, so that the inequality Qvxo< 1 is fulfilled (where xo is the mean length of electron hydration), the expression (12) yields : j = 21when+<-1.2V; j = 2xoQvZwhen~>-0.8V.p was found to be 0.25. By comparing the photocurrent values for these two limiting cases and, thus, determin- ing Q,, it is easy to obtain the rate constant k, which was found to be 5 x lo6 s-I. A value of k, of about the same order of magnitude was found using the photoemission technique with pulsed illumination of the electrode. According to ref. (17), if the NO, concentration is high, emitted “ dry ” electrons can react directly with this scavenger prior to their hydration. Therefore, if it is possible for the “ dry ” electrons to return to the electrode, the photocurrent in the NO, solution should be higher than in solutions of the scavenger that do not capture the “dry” electrons (for instance, H30+).However, this does not influence theYU. v. PLESKOV, z. A. ROTENBERG, V. v. ELETSKY, v. I. LAKOMOV 61 results of the calculations for the rate constant k, since it is determined from the ratio of the slopes for two limiting portions of the (jom4, 4) curve (that is, j/(Z-Ze)) which does not depend on the absolute value of Z-Zc. G. C . Barker, A. W. Gardner and D. C . Sammon, J. Electrochem. SOC., 1966, 113, 1182. Yu. V. Pleskov and 2. A. Rotenberg, Uspekhi Khim., 1972, 41,40. A. Brodsky and Yu. Pleskov, Progress in Surface Science, ed. S . G . Davison (Pergamon, Oxford, 1972), Vol. 2, part 1. 2. A. Rotenberg, V. I. Lakomov and Yu. V. Pleskov, Elektrochim., 1970,6, 515. A. N. Frumkin, Zhur. Fiz. Khim., 1957, 31, 1875. ' S. D. Levina and T. V. Kalish, Dokl. Akad. Nauk S.S.S.R., 1956, 109, 97. ' I. A. Bagotskaya and A. I. Oshe, Trudy 4 Sooeshshaniyapo Electrokhimii (Izdatelstvo Akademii Nauk S.S.S.R., Moscow, 1959), p. 82. L. I. Krishtalik, Uspekhi Khim., 1965, 34, 1831. G. C. Barker, Ber. Bunsenges. Phys. Chem., 1971, 75, 728. B. N. Kabanov, Zhur. Fiz. Khim., 1936, 8, 486. lo L. I. Krishtalik, Elektrochim., 1968, 4, 877. l 2 Z . A. Jofa and K. P. Mikulin, Zhur. Fiz. Khiin., 1944, 18, 137. l 3 Z. A. Rotenberg, V. I. Lakomov and Yu. V. Pleskov, J. Electr-oannl. Clici~i., 1970, 27, 403. Is L. I. Korshunov, Ya. M. Zolotovitsky, V. A. Bendersky and V. 1. Goldansky, Khiin. Yysokikh l 5 V. V. Eletsky, Z. A. Rotenberg and Yu. V. Pleskov, Khim. Vysokikh Energij, 1971, 5, 325. l 7 J . E. Aldrich, M. J. Bronskill, R. K. Wolff and J. W. Hunt, J. Chem. Phys., 1971, 55, 530. Enetyij, 1970, 4, 346. Yu. Ya. Gurevicli and Z. A. Rotenberg, Elektrokhim., 1968, 4, 529, 984.

ISSN:0301-7249    

DOI:10.1039/DC9735600052

出版商:RSC    

年代:1973

数据来源: RSC

摘要:

GENERAL DISCUSSION INTERMEDIATES I" ELECTROCHEMICAL REACTIONS Dr. R. Parsons (Bristol) said : I should like to add to Professor Gerischer's list a new method which may have some prospect of success. This is the application of photoelectron spectroscopy (ESCA) to material adsorbed on electrode surfaces. It is well established' that the strongly bound intermediate in the formic acid oxidation, A 10' COUNTS src-' L I l L h 280 285 290 binding energy/eV of CIS FIG. 1.-Photoelectron spectra of carbon Is line of platinised platinum electrode dipped in ( a ) 0.5 M aqueous sulphuric acid, (6) formic acid, (c) methanol. The electrode was then washed and intro- duced into the Spectrometer which was evacuated without heating the specimen. and the similar intermediate in methanol oxidation on Pt, is retained on an electrode which is removed from the reagent solution, washed and transferred to another solution.Dr. A. Capon and I, together with Drs. G. Allen and P. M. Tucker of C.E.G.B., Berkeley Nuclear Laboratories, have attempted to detect this species using ESCA. The strongly bound intermediate was formed on a platinized Pt electrode which was then washed in de-oxygenated water and then introduced into a Vacuum Generators ESCA spectrometer which was then evacuated without heating ' A. Capon and R. Parsons, J. Elecfroanal. Cllem., 1973, 44, 1, 239 ; 1973, 45, 205. 6'GENERAL DISCUSSION 63 the electrode. The large peak at -285 eV appears on all spectra and is attributed to hydrocarbon material from pump oil. The small peak at -289 eV appears only on electrodes which have been dipped in (a) HCOOH or (c) CH,OH ; when they were dipped in 0.5 M H2S0, (curve (6)) it did not appear.Since the shift in the energy of this peak is comparable to that of carbon attached to oxygen, it seems reasonable to attribute this peak to an intermediate such as the postulated COH adsorbed with three bonds to Pt although the effect of the bonding to Pt is unknown. Although the peak is small these results do suggest that ESCA provides a useful addition to the repertoire of methods for studying intermediates in electrode reactions as well as oxide layers.' However, it must be noted that the species studied is probably not an intermediate in the main oxidation reaction of formic acid although it is in the oxidation of methanol.The resulting spectra for the carbon I s line are shown in fig. 1. Prof. A. A. Week (Prague) said : Prof. Gerischer mentioned in his talk several methods used for the detection and characterization of unstable intermediates of electrode reactions. All methods mentioned are based on the application of physico- chemical techniques for the direct detection of the intermediate. There is, however, an indirect method, usually requiring no special instrumentation but quite a lot of chemical skill and intuition. This method is based, in principle, upon the addition of a suitable substance which reacts very rapidly with the intermediate (I) in a specific and characteristic way, giving a product (P) from the nature of which the nature of the intermediate can be deduced.If the original reaction is described by the scheme electrode Li D=-=*I-+FP, the new reaction sequence is then electrode k i D-=+I-+FP. k 2 i + X P The effectiveness of the method depends upon the specificity of the I + X+P process and upon its rate. The induced reaction has, obviously, to compete with the I-+FP reaction. This can be achieved only if the rates of both processes are at least compar- able. On the other hand, this method can be used in cases where the I+FP process is so rapid that the steady state concentration of I is very low to be detected by available physicochemical methods. Provided a suitable reaction can be found, this extends the possible range of detection of short lived intermediates to those with life-times of the order - lo-' s.Experimentally not the unstable intermcdiates but the, more or less, stable product of the induced reaction (P) is being followed. The detection of this product can be achieved by electrochemical methods or by any other method used to follow the products of electrode reaction. Thus, for example, long term electrolysis in the presence of X is carried out and the reaction mixture is analyzed for the product P. In suitable cases, the product P can be isolated and characterized in more detail. Another means of detecting the induced reaction I+X is, in cases where X is electroactive itself, to follow the changes of the electro- chemical behaviour of X. This approach has been used, for example, for the detection of CH,-anions formed in the course of the reduction of some organo- K.S . Kim, N. Winograd and R. E. Davis, J . Amer. Chem. SOC., 1971, 93, 6296.64 GENERAL DISCUSSION metallics in tetrahydr0furan.l In this case, magnesium bromide (reduced with two electrons) was added to the solution and polarographic curve was recorded. The CH, -anions (or Cryptoanions-CH: OM) are formed at the electrode surface and diffuse from the electrode. Magnesium bromide, diffusing to the electrode, is con- sumed by the reaction forming MgCH,Br, which is reduced by one electron only. The reaction is thus manifested by the decrease of the wave of magnesium bromide. Several other reactions were followed by this method especially in combination with the preparative electrolysis and separation and identification of the product of the induced reaction.Prof. Sir George Porter (Royal Institution) said : Kuwana has described the use of signal averaging for the improvement of signal/noise ratio. Equally effective is to increase the monitoring light intensity which again produces an improvement in signal/noise proportional to the square root of the intensity. Pulsed light sources, which can temporarily increase intensity by a factor of several hundred, are useful for this purpose. Prof. T. Kuwana (Ohio) said : We have used 75-100 W tungsten-iodine lamps in the visible region of the spectrum. For the u.v., we have employed 100 W xenon arc Iainps with highly stabilized power supplied. Although signal-to-noise can be improved by the use of higher intensity, there are two mitigating features. The first is that one must be careful that the light does not intiate any photochemical processes at or near the electrode surface.The second is that a higher lamp wattage does not necessarily produce a corresponding increase in the light intensity through the optical system, since what is important is the energy-density of the image through the entrance and exit slits of the dispersion system. A higher energy lamp may have a greater total radiant energy, but the energy density per surface area of the arc or filament may not increase significantly. We have used lamps with small area filaments or arcs and have operated them near their maximum rated wattages. For measuring small changes in our light intensity (usually less than 1 part in 100) due to some electrochemical perturbation on the cell, we have used a Princeton Applied Research Corp.Model TDH-9 signal averager, or more recently a mini- computer which is locked-in phase with the repetitive waveform driving the potentio- stat. Dr. A. Bewick (University of Southampton) said : In the paper by Kuwana et al. emphasis is placed upon making measurements at very short times in order to evaluate high kinetic rate constants in the decay reactions of intermediates. I wish to point out that in many cases very high rate constants can be obtained from simple steady state measurements if all of the available information is utilised. For example, a rate constant of 7.5 x lo3 1. mol-' s-I is evaluated in our paper at this Discussion from steady state measurements of both the current density and the amount of intermediate present and it illustrates the advantage of combining both electrochemical and optical data.This particular rate constant is well below the limiting value set by the sensiti- vity of the optical technique in determining the steady state amount of material and diffusion controlled rate constants should be accessible. I wish to ask Kuwana about the details of the absorption spectrum he observed for the thianthrene cation radical. The spectrum is not given in his paper and it would be interesting to know if he observed additional peaks due to the formation of A. A. Vlkk and S. Vozka, to be published (Thesis of S . Vozka, (Prague, 1972)).GENERAL DISCUSSION 65 aggregate species similar to those shown on our spectra. This complication needs to be considered when interpreting the kinetic data.Dr. T. Kuwana (Ohio, U.S.A.) said: In general, the larger the rate constant of the homogeneous chemical reaction involving an electrogenerated intermediate, the shorter is the time gate available for optically observing this intermediate following a chronoamperometric step experiment. It is true that for certain mechanisms, a steady state concentration of a reactant may allow the observation time to be moder- ately long even though the rate is fast. For mechanism studies when one does not know the mechanism apriuri, it is best to monitor the optical absorbance over a wide a time duration as possible. The shape of the (absorbance, time) curve can serve then as a diagnostic for a particular mechanism. Dr.B. Kastening (Jiilich) said : The disproportionation mechanism I1 (eqn (14), (1 8) and (19) of the paper by Broman et al.) has not been discussed in terms of a rate- determining electron transfer (eqn (1 s)), which is frequently observed with anion radicals in the presence of water. While this possibility may not be operative for purely aprotic solvents in which the equilibrium constant is strongly in favour of the cation radicals, the presence of water would cause a considerable shift of the equili- brium towards the right and would affect the rate constant kf of the electron transfer. It would be interesting whether this mechanism would result in a working curve (fig. 6) different from curve A and could they account for the experimental results. Dr. T. Kuwana (Ohio, U.S.A.) said: We did not perform any computer simulations of the absorbance-time behaviour during pulse or relaxation using explicit values of kf or kb for reaction (18).We assumed a pre-equilibrium condition and assumed a value of Keg for reaction (18).l* A more detailed spectroelectrochemical study coupled with product analysis is being conducted presently by Dr. Henry Blount of the University of Delaware. At the moment, our experimental data appear to support the working curves (absorbance, time) of either mechanism I or 111, and not mechanism I1 (see fig. 7 of our paper). Dr. W. J. Albery (Oxford) said : For the reaction scheme of eqn (18) and (19) in Kuwana’s paper, kr k b Th*-t+Th*++Th+Th2+ Th2++H,0+Th0 +2H+, k6 the position of the transition state will be determined by the relative values of kb[Th] and k6[H20].The position of the transition state does not depend on the value of kf. However, in deducing the position of the transition state from the observed data an estimate of the value of kf is useful. If kobs = kf then the first step is rate determining, while if kobs 6 kf the first step must be a pre-equilibrium and the second step the rate determining one. Kuwana’s data show that kobs is indeed much smaller than k,; hence the first step must be a pre-equilibrium. H. J. Shine and Y . Murata, J. Amer. Chem. SOC., 1969,91, 1872 ; J. Org. Chem., 1969,34, 3368. V. D. Parker and L. Eberson, J. Amer. Chem. SOC., 1970,92,7488. H. Blount, University of Delaware, private communication, Aug. 1973. 56-C66 GENERAL DISCUSSION Dr. T.Kuwana (Ohio, U.S.A.) said: In re-examining the computer simulation employed for mechanism 11, the disproportionation, we had indeed assumed instant- aneous attainment of equilibrium for reaction (18). The value of Keg = kf/kb was assumed to be small as indicated by ref. (20) and (22) in our paper. Dr. A. Bewick (University of Southampton) said : In reply to Kastening's comment that the disproportionation rate of the thianthrene cation radical might be low, I wish to point out that there are now sufficient data available to set a lower limit to this rate and also to show that the disproportionation mechanism is not tenable. The disproportionation rate must be at least as large as the observed rate of decay of the cation radical which has been attributed to direct reaction with water ( k z 5 1.mol-1 s-l, in our paper and in Broman's paper). Therefore, the lowest value for the disproportionation rate constant, k,, is given by Putting CTh: = lOP3M and CHz0 = IOM gives k, = 5 x lo4 1. mo1-' s-l. of 2.3 x lo-' for the equilibrium constant, K, of the disproportionation A value has been reported by Hammerich and Parker1 from a measurement of the reoersible half wave potential for the oxidation of Th' to Th2+. Although in conjunction with eqn (1) this demands a reasonable value, 1O1O 1. mol-1 s-l, for the rate constant describing reaction between Th2+ and water, it also shows that the disproportionation route is untenable because of the value required for ki. ki = kd/K From we get the impossible value, k i z 2 x 1013 1.mol-1 s-l. Dr. B. R. Eggins (Ulster) said : In view of my comments on the derivation of the dE*/d(pH) relationship, are Albery et al. fully justified in basing their mechanistic conclusions for the order of proton and electron transfers essentially on this relation- ship ? Have they fully considered the possible alternative mechanism of dispropor- tionation of BH. radicals which, as they have clearly shown, occurs in the photo- chemical reaction ? The occurrence of this reaction would depend on (a) a favourable thermodynamic situation and (b) a favourable kinetic situation. (a) clearly exists because of the disproportionation following the photochemical reaction. However, the equilibrium for the disproportionation would also need to be favoured over the second electro- chemical transfer in the case of the proposed ECEC mechanism or over the protona- tion equilibrium for the ECCE and CECE mechanisms.Probably not enough E" and K-values are known to make such comparisons. The case of (b) is even more difficult to resolve. For the ECEC mechanisms the rate of the disproportionation would need to be faster than the heterogeneous rate of the second electron transfer. For the CECE and ECCE mechanisms, one needs to compare homogeneous rates, i.e., between disproportionation and protonation of BH= to BH;. While I am not suggesting that the disproportionation necessarily does occur, there are so many variables that the possibility of its occurrence cannot easily be ruled out. Ole Hammerich and V. D. Parker, Electrochiin.Acta, 1973, 18, 537.GENERAL DISCUSSION 67 Dr. W. J. Albery (Oxford) (partly communicated). Even though BH. is dis- proportionating in the solution, it may not be disproportionating on the electrode, where it can be converted by direct electron and proton transfer to A or CH2. These faster first order processes are possible on the electrode, but not in the solution. The low concentration of BH* makes second order processes, such as disproportionation, less likely than the first order ones on the electrode. In the greater volume of the solution the second order process, being the only one, does destroy the BH*. The experimental results in fig. 3, 4 and 1 1 of our paper1 show that BHo is indeed being either reduced or oxidised with single electron transfers.If the disproportionation mechanism was operative then we might observe no photo current at all and certainly not one with a cathodic and an anodic branch. Dr. R. M. Reeves (Bristol) (communicated) : The lowering of AEs below the calcula- ted diffusion controlled limit for the hydrogen reduction seems similar both qualita- tively and semi-quantitatively to results we obtained recently for the inhibition of the potassium ion reduction by the hydrogen evolution reaction at the dropping mercury ele~trode.~ Although there was considerable uncertainty in the absolute magnitude of the fractional reduction of current, the effect is significant in terms of the measure- ment technique employed. The results are similar to those of Smith and Heint7. but do not suffer from the problems associated with the use of a pool ele~trode.~ Dr.G. C. Barker (A.E.R.E. Harwell) (communicated): In connection with the comment of Reeves it has to be noted that the data in table 1 of our paper refer to a reaction involving hydrogen ions as well as adsorbed hydrogen atoms ments provide no certain information about the kinetics of the correspo involving water molecules, but if it is assumed (improbably) that the transt (in contrast to that for the reaction involving hydrogen ions) is virtualityzero, the average life of an adsorbed hydrogen atom under the conditions prevailing in the experiments of Reeves et aL4 at very negative potentials might be as long as 10-1GO ps. Such a life would not be long enough to produce appreciable surface coverage when hydrogen ion reduction at the DME is entirely diffusion-controlled. An average life for adsorbed H of 100 p s leads to an H atom surface excess of the order of 0.5 pC cm-2 if [H30+] = 5 x M in the bulk of the solution. Howevei, HLFR studies of the hydrogen evolution reaction at mercury have not thus far revealed the gradual decline in the apparent transfer coefficient to be expected if the average life of adsorbed H atoms were to be long at negative potentials when the surface concentration of H30+ is small, but relatively short when the reaction is not entirely diffusion-con- trolled and H atoms are consumed rapidly by the reaction considered in our paper.On the evidence currently available it seems probable that for mercury the surface coverage is slight in the potential range in which the hydrogen evolution reaction usually is studied.I believe, however, that slight steady state coverage, which probably falls progressively as the potential becomes more negative, may well be the cause of some anomolous transfer coefficient values found for the hydrogen evolution reaction at mercury by the faradaic rectification method many years ago.5 W. J. Albery, M. D. Archer, N. J. Field and A. D. Turner, this Discussion. R. M. Reeves, M. Sluyters-Rehbach and J. H. Sluyters, J. Elecfroaiial. Chem., 1972, 36, 287. F. R. Smith and H. Heintze, Canad. J. Chem., 1970, 48, 203. R. M. Reeves, M. Sluyters-Rehbach and J. H. Sluyters, J . Electroanal. Chem., 1972, 36, 287. G. C. Barker in Transactions of the Symposium on Electrode Processes, ed.E. Yeager (J. Wiley, New York, 1961), p. 325.68 GENERAL DISCUSSION Dr. Frank R. Smith (Newfoundland, Canada) said : Concerning the possible slowness of electrochemical desorption of adsorbed hydrogen atoms from mercury cathodes (Barker’s reaction (h), Pleskov’s reaction B’), there is some evidence that desorption may not be faster than H30& discharge from the recent studies of Heintze and Smith1 and of Reeves, Sluyters-Rehbach and Sluyters.2 In our work we sought to measure the yield of HZ, for comparison with the electri- cal measurements by measuring the over-all volume change of the mercury, the aqueous solution and any gas phase developed. Good precision was possible using the cell shown in fig. 1 which enabled increments of 1 pl to be observed for a total catholyte volume of ca.8 ml and a leak rate measured as 0.1 nl s-I. In experiments with this cell at current densities from 40 pA cm-2 to 5 mA cm-2, the yield of H2 ranged from 50 to 79 % of expectation, with a mean of 66.75 % and standard devia- tion of only 3 %. As fig. 2 and 3 indicate, the rate of volume change and hence of H2 evolution was similarly less than expected on the basis of Faraday’s Law. Hydro- gen saturated 0.1 M HCI was used, the reaction and volume change at the anode eliminated from the observations by a closed solution-sealed tap and the apparatus thermostatted at 25.0 & 0.1 “C. Observations were continued long after the current had been interrupted, the invariable result being that there was a H2 deficit of at least 19 % even after 8 x lo4 s.FIG. 1.-Two-compartrnent cell : a, mercury cathode ; b, platinum contact ; c, 0.2 ml pipette, 0.001 ml graduations ; d, gas inlet for saturation with Hz ; e, Teflon-sleeved B10 stopper ; f, solution- sealed stopcock ; g, platinum wire anode ; h, C14 cone in grooved socket. Many years ago Baars and Kayser3 measured volume changes in hydrogen evolu- tion at several metals, principally Pt, Pb, Hg and Cu. With a simple apparatus, deficits from expectation were found, while using a so-called “improved’’ apparatus which, however, contained well-greased taps,4 the volumes measured agreed closely with Faraday’s laws. The only other similar investigations appear to be Jofa’s H-U. Heintze and F. R. Smith, Canad. J. Chem., 1970,48,203. R. M. Reeves, M.Sluyters-Rehbach and J. H. Sluyters, J. Electroanal. Chem., 1972, 36, 101. E. Baars, Marburger Sitzungsber., 1928, 63, 212 ; E. Baars and C. Kayser, 2. Elektrochem., 1930, 36,428. We (Ref. (1) above) also observed fair agreement with expectations when a greased cell was used : normally a Teflon sleeve was used to seal the B10 stopper.GENERAL DISCUSSION 69 400 360 320 280 240 - a 200 s I60 I20 80 40 0 0 2000 4000 6000 8000 10000 12000 14000 tls FIG. 2.-Volume changes during and after electrolysis at 200 p A (crosses, expt. 18), 500 pA (triangles, expt. 16), 1.00 mA (open circles, expt. 37 ; dots, expt. 5), and 2.50 mA (shaded circles, expt. 38). Arrows denote time of cessation of current, or time of potential rise where indicated. Numbered lines indicate theoretical volume changes. Electrode area 0.6 cm-* (expt.5, 16, 18). 0.5 cmz (expt. 37, 38). 200 I75 I50 125 3 t 4 2 loo 75 5 0 25 0 0 12000 24000 36000 48000 60000 72000 84000 tls FIG. 3.-Volume changes during and after electrolysis at 22.5 p A (shaded circles, expt. 43), 50 pA (crosses, expt. 24) and 100 pA (open circles, expt. 30). Arrows denote time of cessation of current. Numbered lines indicate theoretical volume changes. Electrode area 0.6 cmz (expt. 24), 0.5 cmz (expt. 30,43).70 GENERAL DISCUSSION unpublished work on Hg, cited by Frumkin,' Kolotyrkin's experiments with porous lead and Miiller's work3 with a dropping mercury electrode. The Russian work1* apparently involved very low rates of Hz generation and prolonged observations : it is not clear whether the rate of H2 production kept pace with the current but, at sufficiently long times, agreement was obtained. While not regarding our own work4 as definitive, we do believe that there is an element of doubt that hydrogen evolution at mercury (and at lead) is completed in the manner usually assumed, i.e., that electrochemical desorption is much more rapid than H30$ discharge.We suggest that some of the adsorbed hydrogen genera- ted in discharge leaves the mercury surface by diffusion into the bulk metal, later to reappear as H2 gas after the current has been cut off. In experiments with Das,' using highly purified mercury and solution, the slow growth and rapid decay of hydrogen overpotential has been linked to the growth and release of a single H2 bubble from the Hg surface with a period of ca.3 h at 50 pA cnir2. Other evidence that slow processes, such as diffusion in the bulk metal, affect the overpotential is the observation' of a prolonged rise of overpotential at current densities up to 100pA ~ r n - ~ . That all is not as simple as generally assumed is indicated by the inhibition of potassium ion reduction by a product of H30& discharge in aqueous KCI + HCI, deduced6 from impedance measurements on dropping mercury electrodes. My second point concerns the possibility, referred to by Pleskov, that the ionisa- tion of adsorbed hydrogen is an activationless process in some potential region. I would like to suggest a possible means of observing the behaviour of this process directly over a useful potential range.We have demonstrated7 that a potential of +0.2 V against a hydrogen electrode in 0.2 M KOH suffices to ensure rapid ionisa- tion of H atoms permeating through a thin lead foil after cathodic generation on the opposite face. Tf a sufficiently thin foil of lead, or some other high overpotential metal, were used in such a permeation experiment and sufficiently large cathodic H atom generation currents were applied, possibly from a pulsed source, it appears that the question of the activationless nature of Hads ionisation could be settled. If the permeation currents proved to be independent of the potential of the diffusion side of the membrane up to -0.4 or -0.5 V this would be even more convincing than the report of Pleskov et al. Finally, I should like to ask Barker and Pleskov whether they consider that they have satisfactorily reduced the oxygen concentration in their solutions to the point where their results are unaffected by its possible presence.Reactions such as or are possible processes competing with the adsorption of Ha, and with electrochemical desorption, respectively. The first (1) is diffusion-controlled with a rate constant of 2 x lo'* 1. mol-' s-I. Nothing is known about the second process which might also occur coupled with electron-transfer : Hads + O2 + e- (metal) = HOT. (3) ' A. N. Frumkin, Acfa Physicocliint. U.R.S.S., 1943, 18, 23. Ya. M. Kolotyrkin and A. N. Frumkin, Zhur. fiz. Khim., 1941, 15, 346. 0. H. Miiller, in Polarogruphy 2964, ed. G. J. Hills (Interscience-Wiley, N.Y., 1966), Vol.I, p. 319. H-U. Heintz and F. R. Smith, Canad. J. Chem., 1970, 48, 203. S. N. Das and F. R. Smith, unpublished results. ti R. M. Reeves, M. Sluyters-Rehbach and J. H. Sluyters, J. Electroaml. Clieni., 1972,36, 101 . ' I. Caderskg, B. L. Muju and F. R. Smith, Canad. J . Chem., 1970, 48, 1789.GENERAL DISCUSSION 71 Reduction of the HOT ion or of H202 at mercury at potentials negative to -0.5 V against S.C.E. is well known1 with an over-all reaction: HOT +3H,f,+2e- = 2H,O. (4) It is to be noted that the total electron consumption per photo-electron could be 2 if only reaction (3) is involved (the HOT product diffusing away after its formation) but rises to 4 if both ( 3 ) and (4) participate. In conclusion, it is sometimes the practice in aqueous solution radiation chemistry to repeatedly pulse the solution before the actual experiment.This procedure, it is claimed,, lowers [O,] to a level of < 10-lo M. Is there any future for a comparable approach in electron photo- emission studies with electrodes ? Dr. G. C. Barker (A.E.R.E., Harwell) said: Small traces of oxygen usually are unimportant in photocurrent studies because oxygen is reduced by the electrode throughout most of the accessible potential range and, as a result, the oxygen con- centration tends to be vanishingly small in the region where hydrated electrons are captured by hydrogen ions. The reduction product, H202, is an efficient electron scavenger but usually is not present at a sufficiently large concentration to compete effectively for hydrated electrons. Also H202 is unlikely to react more readily with adsorbed H atoms than with free H atoms and, in the case of a deoxygenated solution, thus is not likely to affect significantly the average life of adsorbed H.In practice, in photocurrent work the electron scavenger concentration usually exceeds by several orders of magnitude the concentrations of trace impurities and only in exceptional circumstances do impurities noticeably affect the photocurrent. There might be a case for pre-irradiation of the solutions in work on the residual photocurrent ob- served in the absence of added electron scavengers. Dr. Yu. V. Pleskov, Dr. Z. A. Rotenberg, Dr. V. V. Eletsky and Dr. V. I. Lakomov (Moscow) said : The presence of oxygen traces in solution cannot affect significantly the results of photoemission measurements for the following reasons.These meas- urements are mainly carried out in the potential range in which oxygen is reduced under the limiting current conditions. Since the formation of atomic hydrogen and its diffusion occur at distances comparable with the thermalization length of electrons (20-30 A), i.e., by 3 orders of magnitude less than the mean thickness of the diffusion layer, the oxygen concentration in the " hydrogen " layer is negligible compared with its concentration (a very low one) in the solution bulk. Prof. J. W. Schultze (Germany) said : The evaluation of the experimental data o f the hydrogen desorption includes the assumption that Faraday's law with z = 1 can be applied for the reaction Hg+ H30+ +e--+Hg- H+ H,O (1) Hg+H,,+Hg-H (2) and, correspondingly , reaction can proceed without current. I.M. Kolthoff and J. J. Lingane, Polarograplzy (Interscience, N.Y. 2nd Edn., 1952), Vol. 2, v. 553. E. J. Hart and E. M. Fielden, in Sohated Electron, Advances in Chemistry Series No. 50, 1965, p. 253.72 GENERAL DISCUSSION However, at constant potential e the charge flow of electrosorption reactions is given by the derivative' where y is the electrosorption valency which may be any broken value, in the case of reaction (1) between 0 and 1. An analysis of different systems has shown, that y z 1 in the system Pt/H, but I y l + ( z [ for all cations and anions adsorbed on mercury, e.g., yz0.15 for the adsorption of alkali ions. Hence, in the case of reaction (1) the assumption y = z is a very rough one.One wonders if it can be proved by the experiments described by Barker or Pleskov. Dr. G. C. Barker (A.E.R.E., Harwell) (communicated): The argument of Schultze is quite valid in the context of our non-steady state measurements but it seems unlikely that y for the adsorption of H atoms on mercury will differ appreciably from unity. Even for the adsorption of relatively large alcohol radicals on mercury there is no evi- dence that radical adsorption at constant potential appreciably disturbs the structure of the double layer. In principle, non-steady state photopotential experiments of the type described in our paper should shed light on the parameter mentioned by Schultze but there are many practical considerations which make improbable any such clarification. Steady state photocurrent measurements, of course, tend to be unaffected by the influence of adsorbed intermediates on double layer charge density. Dr.Yu. V. Pleskov, Dr. Z. A. Rotenberg, Dr. V. V. Eletsky and Dr. I. V. Lakomov (Moscow) said : A partial charge transfer during adsorption of atomic hydrogen on mercury or other metals is likely to affect the results of the photoemission measure- ments under non steady-state conditions, when the reactions of electrochemical desorption and ionization occur at a much slower rate than the adsorption rate of atomic hydrogen. To observe this, it is necessary to use pulsed illumination of the electrode with pulse lengths of the order of 10-9-10-10 s. In the case of steady-state measurements, the rates of H adsorption and electrochemical removal are the same and apparently a partial charge transfer cannot be detected.Prof. I. E. Epelboin (Paris) said : Does Dr. Pleskov think that a photoelectro- chemical method could be applied to the study of oxide layers, and particularly to the determination of stoichiometry ? Dr. G. C. Barker (A.E.R.E., Harwell) (communicated): Epelboin may be interested to know that we made some years ago a preliminary study of light-induced electron emission from tantaIum oxide films on tantalum, using films sufficiently thick to avoid emission from the metal into the oxide film. This work proved to be a formidable task as using pulsed light of wavelength 253.7 nm there were large and somewhat irreproducible background currents connected with light-induced film growth that varied peculiarly with potential.We were, however, after much fruitless work, able to detect fairly reproducible changes in photocurrent at constant potential at the more negative potentials when N20 was introduced into or removed from the solution. This current presumbably was caused by the excitation of electrons in valence bands for the oxide layer, followed successively by some evaporation of hot electrons into the solution, the formation of hydrated electrons and the usual steps (except perhaps J. W. Schultze and K. J. Vetter, J. Electroanal. Chem., 1973, 44,63. G. C. Barker and M. J. Williams, unpublished work.GENERAL DISCUSSION 73 the conversion of OH to OH-) that occur when N 2 0 is present. The quality of the results was not such as to stimulate further work along similar lines.Dr. Yu. V. Pleskov, Dr. Z. A. Rotenberg, Dr. V. V. Eletsky and Dr. V. I. Lakomov (Moscow) said : The presence of an oxide layer on the electrode surface can be a source of at least two side effects superimposed on photoemission. These are : (1) the internal photoeffect in the oxide layer (which as a rule has semiconducting properties) and (2) the reduction reaction of the oxide by solvated electrons. Nevertheless, it is to be hoped that the photo-electrochemical methods in general and the photo- emission method in particular afford great possibilities for investigation of oxide layers on electrodes. Dr. D. J. Schiffrin (Argentina) said : It is suggested by Barker et al.‘ that the oxidation of COZ.is a slow process in the potential range - 1.3 to - 1.5 V. The observed decay of AE, (fig. 4) at potentials anodic to - 1.4 V could be equally well caused by a decrease in the value of the reduction rate constant of COT-, or a combina- tion of both effects. Also I would like to ask why the minimum in AE, caused by the diffusion of the OH. radicals (fig. 3), is not observed for the N,O+HCOZ solution. In the latter case, one might consider that the only difference between the two experiments is the replacement on some OH. by CO;. radicals in solution, while keeping the overall radical concentration unaltered. Dr. G. C. Barker (A.E.R.E., Harwell) said: The data in fig. 4 of our paper suggest that the reduction as well as the oxidation of CO, may appear slow when studied at elapsed times smaller than 0.3 p s .As regards the absence of minima in the case of the curves in this figure it has to be noted that, with formate present at high con- centration (5 x M), the OH radicals formed by electron capture are largely and rapidly converted homogeneously to the less reactive intermediate CO,. If at - 1.6 V COT were to be rapidly reduced by the electrode a minimum comparable with those observed in the absence of formate would be observed. Dr. W. J. Albery (Oxford) said : The potential 4 * in Pleskov’s paper is very similar to the potential of zero photo current, E,, in our work2; it is the potential where the photochemically generated intermediate is equally likely to be reduced or oxidised. We obtained information about the mechanism of our reaction from the variation of E* with pH.Pleskov has similar results in his fig. 3. Has he tried to see if his results will fit on to two straight lines, one independent of pH at low values of [H,O+] and one decreasing with [H30+] at higher acidities? Is there any signi- ficance in the line he has drawn in fig. 3? Dr. D. J. Schiffrin (Argentina) said : There is an apparent contradiction between the experimental results presented in fig. 2 and 3 in Pleskov’s paper. The differential coefficients d4/d log (2j0 - j / j ) and d4*/d log [H+] should have the same value if the mechanism of hydrogen removal via the adsorption stage is considered. However, the values are 120 and 60 mV respectively (the latter is the limiting value obtained from the results presented in fig. 3 at high hydrogen ion concentrations). This discrepancy may be caused by several reasons : (a) the order with respect to G. C. Barker, D. McKeowen, M. J. Williams, G. Bottura and V. Concialini, paper at this Discussion. W. J. Albery, M. D. Archer, N. J. Field and A. D. Turner, this Discussion.74 GENERAL DISCUSSION 4 [Hi-] in k20 is 4; a result which would be difficult to interpret ; (6) kinetic complica- tions introduced by the varying nature of the several proton donors present (H,O, HSO, and H30+) at different H30’ concentrations in the K2S04 + H2S04 mixtures, and (c) reaction (C) proceeding both through an adsorption stage and directly from solution. The use of K2S04 as a base electrolyte is not recommended for the study of the properties of He, not only because of the uncertainties regarding the nature of the proton donors, but also due to the complications in establishing the ea; scavenging species. I also feel it is better not to rely heavily on theoretical photoemission expressions to calculate partial currents, i.e., when v < 2, due to the diffuse layer effects contributing to the observed photocurrent,l but rather use a direct comparison at the same potential with scavengers such as N20. Dr. Yu. V. Pleskov, Dr. 2. A. Rotenberg, Dr. V. V. Eletsky and Dr. V. I. Lakomov (Moscow) said : The dependence of $* on pH is described by eqn (4), (5) and (7) if in these equations we put j = j , (i.e., jH = j,) : - + c As is clear from this formula, the (4*, pH) dependence cannot be represented as a simple superposition of two linear functions. In the limiting case, when K’/K$ K”/KX, this dependence becomes linear and has the same slope of 2.3 RT(a+P)F as the dependence of $ on 10g(2j0 -jlj). As follows from fig. 3, at CHJo+ < 1N this case is not realized as yet. The reason is not clear. It may be caused by inappropriate base electrolyte, indeed. Anyhow this question requires further experimental study. + i - - 4 - G. Bomchil, D. J. Schiffrin and J. T. D’Alessio, J. Electround Chem., 1970, 25, 107.

ISSN:0301-7249    

DOI:10.1039/DC9735600062

出版商:RSC    

年代:1973

数据来源: RSC

摘要:

Application of the Photo-electrochemical Effect to the Study of the Electrochemical Properties of Radicals: COT and CHj BY D. J. SCHIFFRIN Instituto Nacional de Tecnologia Industrial, Libertad 1235, Buenos Aires, Argentina Received 1 1 th July, 1973 The photocurrents of mercury in contact with COz, HCOOH and CH,CI solutions are analyzed. The Cog radical anion is oxidized at potentials anodic to - 1.3 V against N.C.E. and is completely reduced in a 1-electron process at sufficiently cathodic potentials. A proton donor is involved in the reduction process and there is a change in the reduction mechanism when going from neutral to acid solutions. The photocurrents for formic acid solutions in the presence of HCOO- ions are inter- preted and the indirect formation of COT is inferred.Methyl radicals are found to be reduced in a 1-electron process at potentials cathodic to - 1.5 V and cannot be oxidized on a mercury electrode. The range of electrode potential stability for these radicals is discussed in relation to electrosynthetic processes. The irradiation of a polarized mercury surface in contact with aqueous electro- lytes leads to the photoemission of electrons into the s o l ~ t i o n . ~ - ~ Since the energy of the emitted electrons is very low (a few eV at the most), their thermalization and subsequent hydration occurs very close to the metal surface,'. and, if an appropriate electron scavenger is present, results in the homogeneous generation of radicals. As the radiation energy required for the photoemission process is much lower than that required for water ionization, the photo-electrochemical effect is a " cleaner " method for generating radical species than are pulse radiolysis techniques.The actual photocurrent measured depends on the decay mechanisms for the radical product of the scavenging process, and a photocurrent is only observed if the radicals decay to electrochemically inert species either through dimerization or reduction at the electrode. In this respect, the photo-electrochemical effect can be used in the study of electrochemical reductions involving the formation of radical inter~nediates,~ 9 since the intermediates can be generated independently of the overall electrochemical process. The purpose of this work is to show the application of the photo-electrochemical effect to the study of some aspects of the reduction mechanism of carbon dioxide and the methyl halides on a mercury electrode, attempting to correlate radiation chemistry results with electrochemical kinetic studies.The first case is undoubtedly of econ- omic relevance and there has been some discussion regarding the reduction mechanism and the products that can be ~btained.~? The second case illustrates the electro- chemical properties of a radical that is not easily oxidized on mercury and which has been postulated as an intermediate in the reduction of the alkyl halides.9 EXPERIMENTAL The experimental arrangement has been described previously 6 - l o and involves the use of a modulated light source and a phase detection technique in the photocurrent nieasure- ments.The circuit 75 The circuit diagram of the drop knocker employed is shown in fig. 1 .76 PHOTO-ELECTROCHEMICAL EFFECT associated with OAl is a square wave generator, the frequency of which could be altered by means of RJ. This signal was differentiated and the resulting pulses clamped and used to drive the pen drop of the X-Y recorder through T1. The same square wave was delayed by approximately 0.1 s, differentiated, clamped, inverted (OA3) and used to drive the drop hammer through T,. The time constant of the differentiating circuits was adjusted to allow for the mechanical inertia of the solenoids operating the pen drop and the hammer. In this way, the photocurrent data could be recorded at a constant time in the life of the drop ; immediately after recording, the drop was knocked off the capillary tip.C. I 1 H - 0 v2 FIG. l.--Circuit diagram of the drop knocker employed. All Op-Amps are Texas Instruments SN 72741. T1 and TZ are 2N 4920 transistors ; H is the solenoid operating the hammer ; B is a 30 V battery pack. C1 = 0.2, C2 = 0.15, C3 = 0.5, C4 = 0.01 and C5 = 2 pF. R1 = 5000, Rz = 12, R3 = 50, R+ = 150, RS = 500, Rs = 1O00, R7 = 100, RB = 47, R g = 10 klR. EIV FIG. 2.-Photocurrents of mercury in contact with a 1 M KCl solution saturated with NzO.D. J. SCHIFFRIN 77 Triply distilled water, twice from alkaline permanganate, was used throughout. HCI was distilled and the azeotropic mixture was employed. A.R. B.D.H. KCI was twice re- crystallized, but A.R. B.D.H. KHC03 and HCOOH, and B.D.H. reagent grade HCOOK and CHJ were used without further purification.The purification of mercury and the prepara- tion of the dropping mercury electrodes has already been described." Before each set of measurements, the photocurrent of a saturated N20 solution in the medium without the scavenger to be studied was recorded. This served as a calibrating procedure for both the light intensity and the drop area. All the results are calculated as relative to the saturated NzO solution under the same experimental conditions and are thus relative values. A typical calibration run is shown in fig. 2. The potentials were measured with respect to a 1 N calomel electrode. The temperature to which all the results refer is 22_+ 1°C ; no attempt was made to use an air thermostat due to the possibility of interference by stray mains noise, but the cell was placed inside a Faraday cage with insulated walls.The Hg lamp was placed outside the cage and the light reached the cell through an orifice in the wall. All the photocurrent values reported here iile in arbitrary units. RESULTS AND DISCUSSION I-CARBON DIOXIDE (a) NEARLY NEUTRAL SOLUTIONS Photocurrent measurements were recorded for 1 M constant ionic strength mixtures of KHC03 and KCl for the concentrations (1 -m)KCl+mKHC03 for rn = 0.05, 0.07, 0.1, 0.2, 0.4, 0.7 and 1 M, and for the same mixtures saturated with 4 .,a 2 r .. .. . . z. 0 .. .. * . ..'."' I I - 1.0 - 1.5 -2.0 EIV FIG. 3.-Effect of deaeration of a 0.05 KHCOJ+0.95 M KCl solution on the photocurrent. (a) No deaeration ; (b) 5 min deaeration, and (c) 8 min deaeration.The instrumental sensitivity in these measurements was 3.16 times greater than the experimental results presented in fig. 4. COz. In all the experiments involving acid carbonate solutions, no previous deaera- tion with nitrogen was performed. Instead, the solutions were prepared with freshly distilled water immediately before use. This procedure was adopted to avoid changes in the equilibrium concentration of COz arising from a shift in the CO;--HCO;- C02 equilibrium. The photocurrents observed with air saturated solutions are at least two orders of magnitude lower than those usually measured in these experi- ments and the trace impurities of oxygen present with the procedure adopted can be safely ignored. The effect of deaeration is shown in fig.3, where the effect of the decrease in equilibrium C02 concentration results in a decrease of the observed photocurrent.78 PHOTO-ELECTROCHEMICAL EFFECT 4 - .,P Fig. 4 shows some of the experimental results for the mixtures. The photocurrent increases with the concentration of KHC03 ; however, the scavenging species is not the HCO; ion since this anion is unreactive towards the hydrated electron (kHcoz < lo6 from its photo-electrochemical activity in its saturated solution in 1 M KCl, where the HCO, ion concentration is very low (fig. 5). The primary product of electron capture is the CO; radical anion,13* l4 and the observed photocurrent will depend on the decay path of this radical. As shown in fig. 4, no photocurrent was detected for potentials anodic to - 1.3 V.This M-l s-l , 12 ). The reacting species is in this case the C02 molecule, as can be seen CO,+ea& -, CO; (1) 2 - EIV FIG. 4.-Photocurrents of mercury in contact with KHC03 + KCl mixtures at 1 M constant ionic strength ; (a) 0.05, (b) 0.1, (c) 0.2, (d) 0.7 M in KHC03. Ot . . ' . . . L ........... -. I I point was also verified by performing experiments with increased instrumental sensitivity and with the full unfiltered output of the mercury lamp. The absence of a photocurrent means that, in this potential range, the CO; radical is completely oxidized at the mercury electrode l 5 and that no homogeneous second order dimer- ization reaction occurs in solution (see below CH3C1 case).D. J. SCHIFFRIN 79 A useful parameter for studying the potential dependence of the photocurrent is the stoichiometric number for the overall process, v, which represents the total number of electrons permanently lost by the electrode per successful electron capture act in so1ution.l.l 5 The value of v at different potentials was calculated from where ip(s) and ip(N20) are the photocurrents observed for the solution of scavenger S and N20, k, and kN20 are the rate constants for the reaction of e& with S and N20, and C, and CN20 are the concentrations of S and N20 respectively. Eqn (2) assumes that the concentration of scavenger is sufficiently low for the limiting behaviour of the photocurrent-scavenger concentration to be valid. The various models for the diffusion problem proposed so far 1* lead to a limiting behaviour of the form where iE is the total photocurrent emitted, D, is the diffusion coefficient of e&, and 6 is an emission distance parameter, the value of which depends on the hydrated electron deposition function chosen.The constant 2 appearing in eqn (2) is the stoichiometric number of the photo-electrochemical reduction of N20, i.e., for each electron captured, an OH- radical is formed which is further reduced to give OH-, and in total, two electrons are transferred to the solution for each successful act of electron capture.1 0.1 0.2 d cc02 FIG. 6.-Plot of the ratio ip(COz)/ip(NzO) against Cco2 for the KHC03 + KC1 mixtures at constant ionic strength, and for the same, saturated with C02, in the potential range - 1.6 to - 1.7 V. In the potential range - 1.6 to - 1.7 V, the photocurrent ratio ip(CO2)/iP(N20) is constant for each C 0 2 concentration.The value of the equilibrium concentration of C02 in the KHC03 +KC1 mixtures was calculated from thermodynamic data l 6 and the solubility of C02 in KCI was taken from tables.17 The test of eqn (2) is shown in fig. 6. An average value of kc02 of 7.95 x lo9 M-l s-l can be estimated for vcoz = 2, in good agreement with pulse radiolysis data (7.7 x lo9 M-l s-').I3 The drop in photocurrent for potentials cathodic to - 1.7 V is due to the depletion of scavenger by its electrochemical reduction, a situation similar to that observed in the80 PHOTO-ELECTROCHEMICAL EFFECT case of H+ and H202 solutions. The value of the stoichiometric number implies that the CO; radical formed as primary product is further reduced in a one electron step reduction, presumably to formate ion, which is the product of the electrochemical reduction of CO, from HC0;-C02 solution^.^ The mechanism for the photo- current must involve then the following reactions : CO,+e& 4 CO; followed by and/or by CO; 3 CO,+e, CO;+BH+eG --+ HCO,+B' I ' t I I I - 0 .8 0 - 1 .oo -I .20 - 1.40 - 1.60 EIV FIG. 7.-Potential dependence of the stoichiometric number for the photo-electrochemical reduction of COz from 1 M constant ionic strength mixtures. (a) 0.05, (b) 0.1, (c) 0.2 and (d)0.7 M in KHC03. Curve (e) corresponds to a 0.7 M KHCOJ+0.3 M KCI saturated with COz. Each curve is shifted upwards one unit of v. where BH is a species that can act as a proton donor in step (5) and eG represents an electron in the metal.The values of the photocurrent and also of vco2 observed at a given potential, will depend on the relative rates of reactions (4) and (5). The potential dependence of vco2 is shown in fig.7. Here the transition from vco2 = 0 to vco2 = 2 in a potential range of 250 mV can be seen to resemble in shape an irrevers- ible polarographic reduction wave. It is also found that vco2 is independent of the bulk concentration of C02. The potential dependence of vco2 can be further analyzed considering the reactions (I), (4) and (5). In this potential range it is considered that reaction (4) and (5) are activated processes. In this case, the rate of generation of HCO; ions is given by l 8 :D. J. SCHIFFRIN 81 where k5 is a constant related to the free energy of activation of the process at a reference potential, as is the transfer coefficient, f i s given by F/RT (F = Faraday, R gas constant and T, the absolute temperature), E is the potential and E; is the standard potential corresponding to reaction (5).Here, the oxidation reaction of HCO; is neglected, since the polarographic oxidation of this ion is very irreversible and cannot be observed on a mercury electrode. If the reverse of eqn (5) could occur to any significant extent, its product, the CO; radical, would be easily oxidized to potentials anodic to - 1.3 V, and an anodic wave should be observed in this potential range for formate solutions. - 1.4 EIV unit of log((2--v)/v) and they are in the same order as shown in fig.7. FIG. 8.-Test of eqn (10) for the same solutions shown in fig. 7. Each curve is shifted upwards one Similarly, the rate of oxidation of COT is given by v4 = k,Ccog exp {(1 --a,) f(E - Ei)). (7) The rate of generation of CO; is given by l and hence u1 = 6QiJF u1 = v4+v, From eqn (2), (6), (7) and (9), At a given concentration of proton donors, eqn (10) predicts a linear relationship between log(2-vco2/vco2) and the potential ; fig. 8 shows this plot for several con- centrations of COz and KHC03 and the linear relationship predicted by eqn (10) is apparent. Furthermore, the values of vco2 are fairly independent of the bulk con- centration of COz, as predicted by eqn (lo), supporting the analysis of the potential dependence of the photocurrent given above.82 PHOTO-ELECTROCHEMICAL EFFECT For vCOr = 1, it is possible to define the equivalent of the half wave potential in polarography.This photo-electrochemical half wave potential, E;, is given by and is the potential at which the rates of reduction and oxidation of CO; are the same.5 Table 1 and fig. 12 show that E$ is almost constant for all the solutions containing KHC03 and varying amounts of COz, and having a pH greater than 6.5. TABLE VALUES OF THE COMPOUND CHARGE TRANSFER COEFFICIENT AND OF THE PHOTO- ELECTROCHEMICAL HALF WAVE POTENTIAL E$IV against N.C.E. compound charge solution transfer coefficient 0.05 M KHC03+0.95 M KCl 1.13 - 1.485 0.07 M KHC03+0.93 M KCl 0.99 - 1.485 0.10 M KHC03+0.90 M KCI 1.03 - 1.485 0.20 M KHCO3+0.80 M KCl 1.03 - 1.475 0.40 M KHC03 + 0.60 M KCI 0.83 - 1.475 0.70 M KHCO, + 0.30 M KCl 0.81 - 1.470 Solutions saturated with COz 0.05 M KHC03+0.95 M KCI 0.07 M KHCO3+O.93 M KCI 0.10 M KHC03+0.90 M KCl 0.20 M KHC03+0.80 M KCl 0.40 M KHCOS + 0.60 M KCI 0.70 M KHC03 + 0.30 M KCi 0.001 M HCI+l M KCI 0.002 M HCI+ 1 M KCI 0.005 M HCI+1 M KCI 0.01 M HCl+ 1 M KCI 1 M KHZP04 1 M NH4CI 0.96 0.84 0.83 0.90 0.83 0.83 0.63 0.61 0.53 0.30 0.57 0.63 - 1.490 - 1.490 - 1.490 - 1.485 - 1.480 - 1.470 - 1.340 - 1.295 - 1.265 - 1.210 - 1.270 - 1.340 From the constancy of E$ it must be concluded that the proton donor involved in reaction ( 5 ) cannot be the HCO, ion or H2C03 acid; hence a water molecule must be involved in the reduction step.Values of the compound charge transfer coefficient 1 +a4+a5 are also shown in table 1, but, unfortunately, it is not possible to obtain more information from eqn (10) and (11) regarding the individual values of the constants appearing in these expressions, although the values of Ei and E,” are related through the standard potential of the CO,/HCO; couple.(b) ACID SOLUTIONS Measurements were performed for 1 M KCl solutions saturated with C02 con- taining 2 x 5 x and 2 x M HCI. Fig. 9 illustrates the behaviour observed. Since two scavengers are present (H+ and C02), the observed photocurrent is a combination of the electrochemical activity of the primary products formed, He atoms and CO; radicals. For these mixtures, it was found that the ratio i,(H++CO,)/i,(H+) is constant at potentials anodic to - - 1.1 V for a given HCl concentration.In this potentialD. J. SCHIFFRIN 83 6 - 4 - a .... 2 - 0 range, it can be considered that vco2 = 0 and vH+ = 2 Is- concentrations studied, this ratio is given by and for the scavenger (b)..*" . 0 . . . -...-. . . - . -. <a) .. ..* *.i . . a . L . * . ... - . ,., ._C ..,.. * ......, ;..;.;.;;;*********' I I t I where I A test of the validity of eqn (12) is shown in fig. 10. The value of kH+ was calculated from the HCI+ 1 M KCI experimental results, with reference to the photocurrents for solutions containing only N20, and assuming that vH+ = 2. A value of 1.4 x 1O1O M-I s-1 was found, which is lower than the rate constant measured in aqueous soIutions (2.36 x 1 O 1 O M-I s-l) l3 extrapolated to zero ionic strength from radiolysis 0 .4 0 .e Q(H+)lQ(H+C02) i,(H++ CO,)/ip(H+) for potentials anodic to - 1.1 V. FIG. 10.-Test of eqn (12) for the competitive scavenging of H+ and CO,. Average values of Dotted line, eqn (12) predictions.84 PHOTO-ELECTROCHEMICAL EFFECT experiments. A lower value for kH+ is in accordance with the Bronsted-Bjerrum theory of saline effects on rate constants of ionic reactions between particles carrying opposite charges, and this effect has already been studied for the H++ei reaction.lg These last results are difficult to extrapolate to the high values of the ionic strength used in the present work. As can be seen in fig. 10, the experimental results are at variance with the predictions of the simple diffusional model employed in eqn (12), in which no allowance is made for the possibility of secondary reactions between the radicals formed and the scaven- gers present.For values of vcoz different from zero, and using N20 as reference, it can be shown that assuming that there are none of the above mentioned complications. The use of eqn (13) to evaluate the potential dependence of vco2 leads to negative values of the stoichiornetric number for C02 at potentials anodic to - - 1.2 V. The failure of eqn (12) and (13) to describe the experimental results can be ascribed to the reactivity of H- atoms with COZY since the apparent rate constant for electron capture by C02 appears to be substantially increased. H- atoms are produced by reaction (14), H++e& -P H. at a rate given by There will be a distribution function of He atom generation which is determined by the photoemitted electron distribution function. The solution of the overall diffusional problem is rather cumbersome, and the simple ideas advanced by Barker will be employed.We assume that, as with the model for electron scavenging, there is a H- atom deposition plane. A fraction of these will return to the electrode, whereas the rest will react with the H- atom scavenger present, in this case, C02, to yield COOH or CO;, H*+CO, -P COOH (164 H*+C02 + COT+H+. (16b) Since H- atom is a powerful reducing agent, it is not unreasonable to expect these reactions to occur. CO; radical-anion is a more strongly reducing agent than its conjugate acidY2O probably the reduction of C02 by H- atoms occurs via reaction 16a, followed by the establishment of the acid-base equilibrium of these radicals.In the potential range studied, all the He atoms reaching the electrode are reduced to H2 through ' 9 ' ' 0 (1 7) He + H+ + ec+ Ifz. The rate of He atom reduction, by analogy with the e& case, is given by where QH. is given by (kH.CcO2/DH.)~. The rate of CO; + COOH generation due to reactions (16a) and (1 6b), is given by 2)16 = ~ Q H * O I Q * (1 9)D. J. SCHIFFRIN 85 From (15), (18) and (19), - 1.4 - 1,s - 1.6 E P FIG. 11 .-Potential dependence of the stoichiometric number for the photo-electrochemical reduction of COa from HCIf KCl solutions ; (a) (b) 2 x (c) 5 x (d) lo-’ and (e) 2 x M in HCI. Each curve is shifted one unit in Y. log c,t FIG. 12.-Hydrogen ion concentration dependence of the photo-electrochemical half wave potential for the photeelectrochemid reduction of COz.86 PHOTO-ELECTROCHEMICAL EFFECT which only differs from eqn (12) by the factor (1 -6QH.).From the results shown in fig. 10, 6QH. = 0.28, and taking 6 N 5 nm 1 * and DH. - cm2 s-', the rate con- stant for the reduction of C 0 2 by He can be estimated as - lo8 M-' s-'. Fig. 11 shows the potential dependence of vcoz calculated using eqn (20) for various HCI concentrations and exhibits the same behaviour as previously observed for the KHC03+C02 mixtures. However, in this acid range, depends on the hydrogen ion concentration, as shown in fig. 12, and the (Eg, log CH+) relationship resembles the pH dependence of the half wave potential for the polarographic reduc- tion of organic compounds,21 indicating the participation of a H+ ion in the rate determining step of the radical reduction, COc+H++ei HCO,.(22) 2 - n # Y h 8 0 - I on I - N 0 W - ' - - 1.30 - 1.20 - 1.40 EIV FIG. 13.-Test of eqn (23) for the photo-electrochemical reduction of C02 from acid solutions ; (a) (6) 2 x (c) 5 x (d) M in HCI. Following the same arguments discussed above for the KHC03 + CO, solutions, the equivalent of eqn (10) is now log r=> = log k4 k,f and Fig. 13 shows the potential dependence of l o g ( 2 - ~ ~ ~ ~ / ~ ~ ~ ~ ) for different HCI concentrations, and the values of the compound charge transfer coefficients are shown in table 1. The experimental results appear to be described by eqn (23), although a change of slope is evident for the more concentrated solutions.The average value of the coefficient 1 -a4+a2, for the M HC1 solutions is - 0.59, in reasonable agreement with the value calculated from the hydrogen ion concentration dependence of EZ (-0.56), which points to the internal self consistency of the experimental results. However, the data in fig. 13 could be equally well described assuming a protonation step preceding the electrochemical reduction, 2 x 10-3 and 5 x CO; +H++COOH (25)D. J. SCHIFFRIN 87 followed by COOH + e;+HCO,. There is strong evidence that the CO; radical-anion is the conjugate base of the carboxylate radical,20* 22-24 and the pK value for the acid-base equilibrium (25) has been estimated at 3.9 22 from transient conductance experiments, 2.8 23 from the pH dependence of radiolysis products and more recently, at 1.4 24 from the pH dependence of the adsorption spectra of the radical.In this case, the equivalent of eqn (23) is K being the equilibrium constant for equilibrium (25). Eqn (27) differs from eqn (23) only in the term -log K and there is no simple way of distinguishing between the two mechanisms since the values of the rate constants and of the charge transfer coefficients are unknown, and both mechanisms predict the same hydrogen ion concentration dependence of the photo-electrochemical half wave potential. We can conclude though, that in acid solutions, the irreversible reduction of CO; involves a H+ ion in the rate determining step. However, it is interesting to speculate further on the pK value of the COOH radical. The possibility of complete protonation of the CO; radical-anion in the acid concentration range studied can be excluded.In this case, the radical reduction pathway should be that indicated by reaction (26), whereas the oxidation process should be This mechanism leads to values of Ez independent of the hydrogen ion concentration and can not be correct. Hence, the reported value of 3.9 for the pK can be considered improbable, although the possibility of the participation of the radical species in the reduction process adsorbed on the mercury surface can not be excluded. The change in the value of the compound charge transfer coefficient for solutions more concentrated than 5 x lo-, M in HC1, is an indication of a change in the re- duction mechanism, probably due to the participation of reaction (26) in the reduction process. The 0.02 M HCl solution results are difficult to understand; the limiting cathodic value of vcoz was found to be significantly lower than 2 (1.52).This might be caused either by secondary kinetic complications of the COOH radical, or by deviations from the limiting behaviour of the photocurrent described by eqn (3). Diffuse layer effects can provide information regarding the nature of the activated state. The effect of the base electrolyte concentration on the photocurrent both in KHCO,+CO, and in KCl+CO, solutions, is shown in fig. 14. When the ionic concentration is altered from 0.1 to 1.1 M, a shift of N - 50 mV in the value of E$ for the case of the KHCO,+CO, mixtures is noted, whereas the photocurrent is almost unaffected in the KCl+CO, case.The shift is in accordance with the idea that a decrease in the value of the outer Helmholtz plane (OPH) potential at constant applied potential will result in an acceleration of the oxidation process and a decrease in the reduction rate of a negatively charged species, and can be taken as additional confirmation for the proposed reduction mechanism of the CO; radical-anion. It is not easy to understand the invariance of E-: for the KCl solutions (acid range) on COOH+CO, + H+ + eG. (28)88 PHOTO-ELECTROCHEMICAL EFFECT the assumption of a pre-protonation step mechanism, unless the rate of reaction (4) is independent of the potential in the potentiai range con~idered.~~ If the CO; re- duction occurs through reaction (22), the invariance of Eg is caused by a compensation of the effect of the OHP potential variation, on the oxidation and reduction rates, and the local changes in the hydrogen ion concentration at the electrode surface. 6 4 *! 2 0 - I-;o - 1.5 - 2.0 EIV FIG.14.-Diffuse layer effects on the photocurrent of mercury in contact with C02 saturated solutions of V, 0.1 M KHCO, ; El, 0.1 M KHC03+1 M KCl; 0 , O . l M KCl and 0, 1 M KCI. (C) THE EFFECT OF PROTON DONORS The nature of the proton donor in the reduction step was studied using photo- current measurements in C 0 2 saturated solution of 1 M H2Na2P04 and and 1 M NH4Cl. The E? values obtained are more anodic than the values corresponding to the hydrogen ion concentration in solution, indicating that the H2P0, and NHf; ions are participating in the reduction step as proton donors, with an order of strength of H2P0, >NHt > H20, in accordance with their pK values.(d) THE ELECTROCHEMICAL REDUCTION OF COz Carbon dioxide is electrochemically reduced to formic acid in aqueous solutions ; the intermediate formation of adsorbed CO; and COOH adsorbed radicals has been po~tulated.~ From the photocurrent experiments, no evidence can be found for a radiation induced coverage of the electrode surface with radicals at potentials anodic to - 1.4 V and there is no limiting coverage potential region where the electrochemical reduction process either does not occur or it occurs extremely slowly. In fact, the COG or COOH radicals appear to be extremely unstable with respect to their oxidation or reduction when they are generated photo-electrochemically.Furthermore, some electrode admittance measurements using phase detection techniques 26 failed to reveal any adsorption peak in the electrochemical reduction of C 0 2 from 1 M KCl andD. J. SCHIFFRIN 89 1 M (CH3)4NCl solutions. For the latter case, it has been suggested that the hydro- phobic character of the layer of adsorbed alkyl cations favours the adsorption of the CO; radical-anion on the electrode surface, leading to the possibility of complex radical reactions. No evidence for this could be found in a photo-electrochemical experiment in this medium (fig. 15). The photocurrent potential dependence is identical to that observed for 1 M KCl, but shifted 70 mV cathodic. Also, a second wave is observed at potentials cathodic to - 1.65 V, which is caused by the onset of the photo-electrochemical reduction of (CH314N+ (wave (a) in fig.15). It can be con- cluded, then, that any electrosynthetic process occurring in aqueous solution at 6 0 - 0 , s -1.0 - 1.5 - 2.0 EIV FIG. 15.-Photocurrents of mercury in contact with (a) 1 M (CH,),NCl; (b) 1 M (CH,),NCI saturated with C 0 2 and (c) 1 M KC1 saturated with C 0 2 . potentials anodic to - 1.3 V and involving the intermediate formation of CO;, will result in the generation of C02. It also seems unlikely that a product other than HC02H could be obtained from the solutions studied in the present work, at po- tentials cathodic to - l .6 V, where the radical-anion reduction is diffusion controlled, unless an alternative to reaction paths (5) and (22) is provided.11-FORMIC A C I D It has been proposed that the primary product of the reaction of a hydrated electron with a carboxylic acid decomposes in two parallel first order processes according to either 27-29 RCOOH’+RCOO-+H* (29) RCOOH-+RCO*+OH-. (30) H* + HCOO-+H2 + CO;. (31) Ha atoms are known to react with formate ion 309 31 to yield CO; radicals through Since the reactivity of He atoms towards formic acid is much less than with the an- 33 it would be expected that in a photo-electrochemical experiment when90 PHOTO-ELECTROCHEMICAL EFFECT 6 - .,a 4 HCOO- ion is in sufficient excess, the He atoms generated through reaction (29) would result in the formation of an equivalent amount of CO; radicals. As a consequence, the photo-electrochemical properties of formic acid should resemble to some extent those of the C02 solutions, besides the complications introduced by the reduction or oxidation of the radical produced through reaction (30).To test this point, the photo- currents of mercury in contact with solutions of HCOOH in 1 M HCOOK were measured. The concentrations studied were 0.05, 0.1, 0.2, 0.4, 0.7 and 1 M in HCOOH ; some results are shown in fig. 16. The photocurrent variations with - -0.5 - 1.0 - 1.5 -2.0 EIV FIG. 16.-Photocurrents of mercury in contact with solutions of (a) 0.1 M HCOOH+ 1 M HCOOK and (6) 1 M HCOOH+l M HCOOK. potential are more complex than for the C02 case; the value of the photocurrent is low up to a critical potential, and a sharp increase is noted at potentials cathodic to this.At potentials anodic to the pre- wave, a small photocurrent is observed, which increases with increasing formic acid concentration. This photocurrent probably originates from the reduction of He atoms to H2 through the participation of He in the electrochemical desorption step of the hydrogen evolution reaction (reaction (17)). This current is proportional to the fraction of the He atoms generated through reaction (29) which does not react with HCOO- ions. There is also a small contribution arising from the reaction of ea; with equilibrium Hf ions coming from the dissociation of formic acid. There is a narrow potential range for both waves in which the radical reduction process is diffusion controlled, i.e., the ratio i,(HCOOH + HCOOK)/i,(N,O) is independent of potential.The concentraion dependence of this ratio in the diffusion controlled regions is shown in fig. 17 and follows the simple diffusional models proposed.'' The rate of reaction (29) is known to be slower than (30), since only -30 "/o of electrons scavenged are transformed into He atoms.29 Thus the pre-wave is due to the reduction of CO; formed through reaction (31), and the main wave due to the reduction of the formyl radical CHO., or its hydration In principle, it should be possible to calculate the value of the fraction a of the radicaI HCOOH- decomposing according to reaction (29) and the stoichiometric Also, a pre-wave appears before this point.D. J. SCHIFFRIN 91 number for reaction pathway (30) from the ratio of wave heights and the formic acid concentration dependence of the photocurrent.From eqn (3), (32) 2a ~- ip( 1”wave) - - i , ( ~ ~ w a v e ) ~ ~ ~ ( 1 -a) and i,(HCOOH)/ip(N20)] dCAcoo,, t = +(*;) (2a+v,,(l-a)). (33) Z’(cHC00H) FIG. 17.-Formic acid concentration dependence of the ratio ip(HCOOH + HCOOK)/i,(N,O) for (a) pre-wave and (b) main reduction wave in the photo-electrochemical reduction of formic acid. - 1.10 - I.3? r1.50 -1.00 - 1.20 - 1.40 EIV EIV FIG. 18.-(a) Potential dependence of vZ9 in the photo-electrochemical reduction of formic acid in 1 M HCOOK. (1) 0.1, (2) 0.2, (3) 0.4, (4) 0.7 and (5) 1 M HCOOH. (6) Potential dependence of log((2-v)/v) for the same solutions. The curves are displaced upwards half a unit with respect to each other.92 PHOTO-ELECTROCHEMICAL EFFECT Using the value of kHCmH of 1.4 x lo8 M-' s-l , l 3 values of a = 0.42 and v ~ ~ - 5 are obtained. It is difficult to rationalize the large value of v, a result that is probably due to the value of k,,,,, used in the calculation being too low.The pre-wave was further studied under conditions of increased instrumental sensitivity. The stoichiometric number for the reduction of CO; was calculated by subtracting from the ratio i,(HCOOH + HCOOK)/ip(N20), the constant value of this ratio at potentials anodic to the wave and due to He reduction, and taking the 0 - a5 -100 log CHCOOH FIG. 19.-Formic acid concentration dependence of E$ for the first photo-electrochemical reduction wave of HCOOH in 1 M HCOOK. value of vcoz = 2 at the end of the wave.These results and also the potential depend- ence of 10g(2-v2,/v2,) are shown in fig. 18, and the average value of the compound transfer coefficient calculated from these data is 0.54 suggesting a reduction mech- anism similar to that of KH2P04 and NH4Cl solutions. Again in this case, the potential dependence of the stoichiometric number, is described by eqn (10). This result was also confirmed by the formic acid concentration dependence of E;-* (fig. 19), from which a value of the compound transfer coefficient of 0.58 can be calculated, in good agreement with the previous result. 111-METHYL HALIDES The electrochemical reduction of CH3Cl is a very irreversible process 9 9 and it has been suggested that the rate determining step is the halogen-carbon cleavage, The CH3- radical formed is reduced to CH4, CH3Cl+e+CH3-+ C1-.followed by CH3*+e;-+CH; CH ; + H 2 0 4 CH4 + OH-. 35, 36 bond (34) (3 5 ) (36) It is possible to by-pass the irreversible step (34) by generating methyl radicals homogeneously by the reaction of photo-electrochemically generated hydrated electrons with CH3Cl. This reaction has been studied with pulse radiolysis tech- nique~.~' Fig. 20 shows the photocurrents observed with this scavenger. There is a sudden change in the potential dependence of the photocunent between - 1.4 and - 1.6 V, and the ratio ip(CH3Cl)/ip(N20) increases by a factor of two. It is suggested that this is due to the occurrence of reaction (35) ; hence the value of v~~~~~ changes from 1 to 2 in this potential range. The potential dependence of the stoichiometric number is shown in fig.21 ; the results resemble a reversible polarographic wave.D. J. SCHIFFRIN 93 I I I I I - 0 . 5 - 1.0 -1.5 -2.0 E/V FIG. 20.-Photocurrents of mercury in contact with a 1 M KCI solution saturated with methyl chloride. By comparing the photo-electrochemical behaviour of COz with CH,Cl, it can be concluded that the CH3- radical is not oxidized on a mercury electrode at potentials anodic to - 1.4 V, and the decay mechanism must be their second order recombina- tion. As a consequence, for each electron capture act, a stable compound is finally formed in solution (C,H,). In the potential range - 1.4> E> - 1.6 V, the observed photocurrent will have two contributions, where i(CH,) and i(CH:) are the partial currents due to CH3- and to its reduction product.It it is assumed that process (35) is reversible, the surface concentration of these radicals is given by the Nerst equation, i,(CH,Cl) = i(CH3*) + i(CH3) (37) -1.00 -1.20 -t.40 -i.60 -1.00 FIG. 21 .-Potential dependence of the stoichiometric number for the photo-electrochemical reduction EIV of CH3CI. Solid line : values calculated from eqn (39) with n = 1 and ELf = - 3.47 V.94 PHOTO-ELECTROCHEMICAL EFFECT where n is the number of electrons involved in the reduction step. Since the radicals are formed at a very short distance from the mercury surface (- 50 A),' the ratio of the partial currents due to CH,. and CH3 (or its decay product, CH4), will be given by 8. From (37) and (38), - 1.40 - 1.50 - 1,60 EIV FIG.22.-Test of eqn (39) for the photo-electrochemical reduction of CH3CI. In fig. 21, the values of v ~ ~ ~ ~ , calculated from eqn (39) for n = 1, are found to be in reasonable accord with the experimental results. However, a more critical test of eqn (39) is shown in fig. 22 ; the value of the slope in this graph is (50 mV)-l instead of the theoretical value of (58 mV)-l. The value of n must equal 1 since vCHsCl changes from 1 to 2, and some degree of irreversibility for reaction (35) would be apparent, although there is too much experimental uncertainty to establish this firmly. The photo-electrochemical half wave potential is - 1.47 V. However, the use of eqn (37) would not be valid if the protonation rate of CH3 radical anion is much faster than its generation rate.It is not possible at present to decide between these two possibilities. Nevertheless, since ET is anodic to the polarographic E+ for CH3Cl (E+ = -2.27 V in a dioxan-water mixture 9), this can be taken as confirmation that reactions (35) and (36) are not rate determining in the Jectrochemical reduction of CH3Cl. Furthermore, the value of E$ gives the anodic limit for the reduction poten- tials of the methyl monohalides to yield CH4. From an efectrosynthetic point of view, this result might imply that any process occurring on mercury at potentials cathodic to - 1.55 V in aqueous solutions, involving the intermediate formation of the methyl radical, should necessarily result in the formation of CH4, and no radical addition reaction can be observed unless reaction (35) is blocked or an alterncttive reaction pathway to reaction (36) is provided. It is interesting to speculate further on the physical meaning of the value of EZ found.The real potential of CH;,,,, can be estimated, if it is assumed that the reduction process is that indicated by reaction (39, and the methyl radical is not adsorbed on mercury. Data for the real potential of H+, the ionization energy of H(,), the dissociation energy of H2(,,, the free energy of solvation of CH;,,, and the electron affinity of CH;,,, were taken from the literature,16* 3 8 9 3 5 9 39 and a value of ctcH3~ of -229 kJ mol-1 was found. The radius of CH; can be estimated to be - 1.14 from the covalent radii of C and H,38 although the presence of an extraD. J.SCHIFFRIN 95 electron in the CH3. molecule could be expected to result in a somewhat larger value of the radius. The value of cx,,~ calculated is much lower than that of anions having even larger radii, for instance, aF- = -446 kJ mol-1*16 implying that the CH3* radicals participating in reaction (35) must adsorb strongly on mercury. More experimental work using pulsed light is required to elucidate this point. CONCLUSIONS The photo-electrochemical effect can give additional information in electrochemical reduction processes involving the formation of a radical intermediate species. The main advantage of this method in kinetic studies is the possibility of observing only the electrochemical activity of radicals generated independently of the reduction process.G. C. Barker, A. W. Gardner and D. C. Sammon, J. Electrochem. SOC., 1966, 113, 1182. Yu. V. Pleskov and Z. A. Rotenberg, J. Electroanal. Chem., 1969, 20, 1. P. Delahay and V. S. Srinivisan, J. Phys. Chem., 1966, 70,420. G . Bomchil, D. J. Schiffrin and J. T. D'Alessio, J. Electroanal. Chem., 1970, 25, 107. Z. A. Rotenberg, V. I. Lakomov and Yu. V. Pleskov, J. Electroanal. Chem., 1970, 27, 403. D. J. Schiffrin, Croatica Chem. Acta, 1972, 44, 139. W. Paik, T. N. Andersen and H. Eyring, Electrochim. Acta, 1969, 14, 1217. A. Bewick and G. P. Greener, Tetrahedron Letters, 1969, 53, 4623. M. v. Stackelberg and W. W. Stracke, 2. Elektrochem., 1949, 53, 118. l o R. de Levie and J. C. Kreuser, J. Electroanal. Chem., 1969, 21, 221. D. J. Schiffrin, Trans. Faraday Soc., 1971, 67, 3318. l 2 J. K. Thomas, S. Gordon and E. J. Hart, J. Phys. Chem., 1964, 68, 1524. l 3 S. Gordon, E. J. Hart, M. S. Matheson, J. Rabani and J. K. Thomas, Disc. Faraday Suc., 1963, l4 J. J. Weiss, Radiation Res. Suppl., 1964, 4, 141. l 5 G. C. Barker, Electrochim. Acta, 1968, 13, 1221. l 6 R. Parsons, Handbook of Electrochemical Constants (Butterworth, London, 1959). 36, 193. International Critical Tables (McGraw Hill, New York and London, 1928). P. Delahay, New Instrumental Methods in Electrochemistry (Interscience, New York, 1966). l 9 E. J. Hart and M. Anbar, The Hydrated Electron (Wiley-Interscience, New York, 1970). 2o A. L. J. Beckwith and R. 0. C. Norman, J. Chem. SOC. B, 1969,400. 21 P. Zuman, D. Barnes and A. RyvolovCKejharova, Disc. Faraday Soc., 1968, 45,202. 22 A. Fojtik, G. Czapski and A. Henglein, J. Phys. Chem., 1970, 74, 3204. 23 J. P. Keene, Y. Raef and A. J. Swallow, in Pulse Radiolysis, ed. M. Ebert, J. P. Keene, A. J. Swallow and J. H. Baxendale (Academic Press, London, 1965). 24 G. V. Buxton and R. M. Sellers, J.C.S. Faraday I, 1973, 69, 555. 2 5 L. I. Krishtalik, Advances in Electrochemistry and Electrochemical Engineering, ed. P. Delahay 26 A. Bewick, K. Kratochvil and D. J. Schiffrin, unpublished results. "J. K. Thomas, Radiation Res. Suppl., 1964, 4, 87. 29 J. K. Thomas, in Advances in Radiation Chemistry, ed. M. Burton and J. L. Magee (Wiley- 30 E. Hayon and M. Moreau, J. Chim. Phys., 1965, 62, 391. 31 J. H. Baxendale and P. L. T. Bevan, in The Chemistry oflonization and Excitation, ed. G. R. A. 32 J. H. Baxendale and D. Smithies, 2. phys. Chem. (Frankfurt), 1956, 7, 242. 33 J. Rabani, J. Amer. Chem. SOC., 1962, 84, 868. 34 A. J. Swallow, Photochem. Photobiol., 1968, 7, 683. 35 N. S. Hush, 2. Elektrochem., 1957, 61, 734. 36 M. G. Evans and N. S. Hush, J. Chim. Phys., 1952, 49, 159. 37 T. I. Balkas, J. H. Fendler and H. Schuler, J. Phys. Chem., 1970, 74, 4497. 38 E. Cartmell and G. W. Fowles, Valency and MoZecuZar Structure (Butterworths, London, 3rd 39 F. Gaines and F. M. Page, Znt. J. Mass Spectrom. Zun Phys., 1968, 1, 315. (Wiley-Interscience, New York and London, 1970), Vol. 7. J. H. Baxendale, Radiation Res. Suppl., 1964, 4, 114. Interscience, New York, 1969), Vol. 1. Johnson and G. Scholes (Taylor and Francis, 1967). ed., 1966).

ISSN:0301-7249    

DOI:10.1039/DC9735600075

出版商:RSC    

年代:1973

数据来源: RSC