Electrochemistry, Past and Present - ACS Publications - American

Chapter 11

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Historical Development of the Understanding of Charge-Transfer Processes in Electrochemistry B. E. Conway Chemistry Department, University of Ottawa, Ottawa, Ontario K 1 N 6N5, Canada The true nature of electrolytic processes in electrochemistry took many years to be understood. An historical outline of the development of these ideas from the pre-Faraday period until the present time is given. One of the matters of outstanding importance for chemistry and electrochemistry was the eventual realization that electricity itself is "atomic" in nature, with the electron as the natural unit of electric charge. Not until this concept was established experimentally, and understood in its theoretical ramifications, was it possible for the microscopic basis of electrolytic processes to be established, and developed more quantitatively with the correct qualitative basis. The final and correct perception of the nature of these processes provided one of the important bases for recognition of the electrical nature of matter itself and the foundations of physical chemistry. In this paper, an historical outline is given of the principal developments in electrochemistry concerning the mechanism and phenomenology of charge transfer in electrolytic processes. In tracing early contributions in this topic, it will be necessary to examine first the historical evolution of ideas about electricity and electric charge. Concepts of Electricity, Charge and the Electron Historically, the understanding of charge-transfer processes in electrochemistry required a long period of gestation. It depended on and, in fact, led in part to, recognition of several complementary factors: the nature of electric charge itself; the association of electric charges with atomic and molecular species in chemistry, giving rise to ions; the solvation and movement of such ions in liquids and some solids, giving rise to conductivity; the discharge or formation of such ions at metal interfaces, characterizing the process of electrolysis and perception of the role of electrons as

0097-6156/89/0390-0152$06.00/0 © 1989 American Chemical Society

In Electrochemistry, Past and Present; Stock, J., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 1989.


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a delocalized system of mobile charges in metals with an energy distribution represented by an energy "band structure". The understanding of the nature of electric charge and, in particular, the atomic entity of electricity, the electron, played a principal role in the foundations of physical and electrochemistry in the last century, and in recognition of the electrical nature of matter. The latter aspect was not fully understood until the earlier years of this century through the development of electronic theories of valency and atomic and molecular structure, chemical bonding and spectroscopy, and electronic theories of metals. In these developments, especially in the last century, electrochemistry played a major role, especially through Faraday's investigations on exchange of charge in processes at metal interfaces that we now call electrolysis, a term, like many others in electrochemistry, that was first used by Faraday himself. It must first be pointed out that the idea and phenomenology of electric charge originated in the 18th century work on the physics of electrostatics, well before any of the principal developments of electrochemistry historically associated with the names of Luigi Galvani, Alessandro Volta, Sir Humphry Davy and Michael Faraday. The phenomenology of generation, storage and transfer of electricity as electric charge by use of electrostatic machines (the Wimshurst machine) and other devices (the Leyden jar and the Electrophorus) was well developed by the middle of the 18th century. (One of the most impressive collections of historically significant examples of such machines is to be found in the Museum of the History of Science in Florence, Italy). However, the "nature of electricity" remained for many years later, through Faraday's time, essentially a mystery and a source of active controversy(1). Electricity was regarded, somewhat like the Aether, as an imponderable fluid and its intimate connection with the structure and properties of matter itself was not appreciated until later in the mid 19th century. In fact, the connection between electrostatic and electrochemical, as well as electromagnetic, manifestations of electricity took a very long time to crystallize; indeed, for a lengthy period, two "kinds" of electricity continued to be referred to: electrostatic and current or Galvanic electricity. The stimulation of the frog's leg nerve in the historical experiment of Galvani, with electricity generated from the contact of two dissimilar metals, marked the point of origin of electrochemistry and its divergence from electrostatics. At first, and indeed for some length of time, the so-called electrostatic (or "common") and "Galvanic" or "animal" manifestation of electricity were thought to correspond to two distinct types of electricity. Again, it was a relatively long time between the first demonstration of the Galvanic effect (1786) and its controlled utilization in Volta's pile (1800) for production of a steady voltage and a current of electricity. During that period, much controversy raged about the supposed differences of the "two kinds" of electricities, and later about electromagnetically induced currents of electricity(1). It could be said, therefore, that until the basic nature of electricity was recognized, the electrochemistry of charge-transfer processes could not be understood and taken further. However, historically, this would be an incorrect perception as the study of



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the nature and phenomenology of charge-transfer processes in electrolysis and in conductivity of salt and acid solutions, was itself precisely one of the main bases of the ultimate understanding of the electrical nature of matter and of electricity itself. The second and complementary basis was, of course, the discovery and characterization of the electron and its charge(e)-to-mass(m) ratio by J.J. Thomson(2), and its charge in experiments by Townsend(3) and by Millikan(4). These experiments demonstrated the ubiquitous involvement of a negatively charged particle, termed the "electron" by Johnstone Stoney(5-6), in electrically stimulated ionization of gases. Electron photoemission from metals was another key discovery. Interestingly enough, for chemistry, one of the most thorough accounts of static electrical phenomena was that published by Joseph Priestley, "The History and Present State of Electricity", London, 1767, in two volumes. It was an article in Encyclopaedia Britannica on the history of electricity by Tytler, copiously illustrated by diagrams of electrostatic machines, based in part on Priestley's volumes, that first stimulated Faraday to construct his first scientific instruments. Generation of static electricity was based primarily on the use of friction machines and the effect became later widely known as "triboelectricity". In the mid-18th century, it was not, of course, recognized that static electricity originated from a mechanical separation of electronic charges (i.e. ionization) at suitable interfaces, classically amber, or glass and resins. However, it is interesting that in one of the earliest lectures given by Faraday he took up a "two-fluid" theory of Major Eeles (1771) and described electricity as a compound body capable of being divided by friction or excitation into two "Portions or Powers" one of which, when separated, always attracts itself to the rubber; the other to the excited Electric. It was also recognized that the "separated" electricities would tend spontaneously to reunite, but the same electricities to repel one another. Faraday was also influenced, at an early age, by a book by Jane Marcet,"Conversations on Chemistry"(2 volumes, London, 1809), written for an audience newly created by Sir Humphry Davy's Royal Institution lectures. In these two volumes, Faraday was introduced to electrochemistry and, according to the author, electrochemical effects, like the origin of static electricity, seemed to require two "electrical fluids", i.e. states we now recognize as positive and negative charge associated with excess or deficiency of electron density. At this point a quotation from Mrs. Marcet's book must be given as it is of much historical significance for chemistry and electrochemistry in relation to the role of electrical effects: "Mr. Davy...whose important discoveries have opened such improved views on chemistry, has suggested an hypothesis which may throw great light upon that science. He supposes that there are two kinds of electricity with one or other of which all bodies are united. These we distinguish by the names of positive and negative electricity; those bodies are disposed to combine which possess opposite electricities, as they are brought together by the attraction which these electricities have for each other. But



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whether this hypothesis be altogether founded on truth or not, it is impossible to question the great influence of electricity in chemical combinations." It is clear from this paragraph that a close connection between chemistry and electricity was already perceived in Davy's time through his work on applications of electrolysis, including the preparation of several of the most electropositive and electronegative elements for the first time. However, it remained for Faraday to quantify the relations between passage of current (charge) and chemical change as expressed in his two laws of electrolytic action: "....the chemical power of a current of electricity is in direct proportion to the absolute quantity of electricity which passes; and "....the equivalent weights of bodies are simply those quantities of them which contain equal quantities of electricity" Thus were the essentials of chemical combination first enunciated! Faraday's own commentary on the consequences of the laws of electrolysis is of interest: "The harmony which this theory of the definite evolution and the equivalent definite action of electricity introduces into the associated theories of definite proportions and electro-chemical affinity, is very great. According to it, the equivalent weights of bodies are simply those quantities of them which contain equal quantities of electricity, or have naturally equal electric powers; it being the ELECTRICITY which determines the equivalent number, because it determines the combining force. Or, if we adopt the atomic theory or phraseology, then the atoms of bodies which are equivalents to each other in their ordinary chemical action, have equal quantities of electricity naturally associated with them." This is an interesting statement, historically, not only for electrochemistry but chemistry in general, for it indicates that Faraday almost, but not quite, reached the vital and important conclusion for chemistry and physics that electricity itself had an "atomic" nature, the electron, as we now know it. Despite his own conclusion cited above, Faraday tended to adhere to older ideas of electricity as a fluid. The "atomic" view of electricity seems not to have been clearly stated, however, until it was expressed in von Helmholtz's Faraday Memorial Lecture (1) in 1881, although a similar idea was advanced by G. Johnstone Stoney at the British Association for Advancement of Science in 1874(5) but not published until 1881. In his later paper of 1891(6), the name "electron" was proposed for this unit of charge, but the latter was not associated with a particular, physically significant, "particle" until Thomson's classic experiment(2) (1897) in which the charge to mass ratio, e/m for the electron "particle" was definitively measured. (Shuster(7) had previously made an electromagnetic deflection experiment similar to Thomson's but found values about 500 times too large). Probably the first to recognize the possibility that atoms, as ions, might be associated with a definite fundamental unit of



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electric charge was Clerk Maxwell in 1873(8). With some reservation, he stated the following with reference to the electrical properties of electrolytes: "If we....assume that the molecules of the ions within the electrolyte are actually charged with definite quantities of electricity, positive and negative, so the electrolytic current is simply a current of convection, we find that this tempting hypothesis leads us into very difficult ground. Suppose we leap over this difficulty by simply arresting the constant value for the molecular charge, and that we call this....one molecule of electricity." One of the difficulties in understanding the true mechanism of the process of electrolysis at the time that Faraday enunciated his laws of electrolysis was the absence of the idea of spontaneous electrolytic dissociation, postulated much later by Clausius and by Arrhenius. In fact, Faraday believed that the electric force at (between) electrodes split up molecules in the electrolyte, giving rise to conductivity. This idea was connected with Freiherr von Grotthus's theory of a series of dissociations and recombinations of charged species in the conductance of aqueous solutions. We now recognize(9) that two significant physico-chemical concepts that had not been developed at that time, gave rise to these conceptual difficulties: one was that it was not realized that socalled salts were already ionic in the pure solid state and the other was that it was not understood that a large solvation (hydration) energy would accompany dissolution of ions of a salt in a polar medium such as water, and largely compensate for the attractive energy amongst ions of unlike sign of charge in a crystal lattice. In fact, an elementary theory of ionic hydration, in terms of dielectric polarization, was not available until Born's treatment in 1920(10) or, in more molecular terms, until Bernal and Fowler's classic paper in 1933 (see also ref. 10). We thus see that a surprizingly long time elapsed between the formulation of the Faraday Laws (1834) and a proper understanding of their significance for chemistry and electrochemistry. Also, until an electron theory of metals was available, no proper microscopic description of charge-transfer in electrolytic processes at electrode interfaces was conceivable. Also, the theory of ionic dissociation in solution in relation to electrolytic conductivity was only developed slowly and accepted with difficulty and much contro-versy through the works of Arrhenius, van't Hoff and Ostwald (the "Ionists"--see the paper in this symposium by K. J. Laidler); in fact, it was not until the early 20th century that "complete dissociation" of so-called "strong" electrolytes became accepted through the works of Lewis, Bjerrum, and Debye and Hückel. Treatments of Charge-transfer in Electrode Processes Historically, in Nicholson and Carlisle's demonstration (April, 1800) of water electrolysis, producing hydrogen and oxygen, the cathodic hydrogen evolution reaction, with proton discharge as its initial step, was the first electrolytic decomposition reaction to be effected following development of Volta's pile, utilizing electricity generated by electrochemical action at dissimilar metals wetted by an



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electrolyte solution. Nicholson and Carlisle's experiment resulted from a personal communication to them from the then President of the Royal Society, Sir Joseph Banks, to whom Volta had earlier commun icated a letter for publication which appeared in the Philosophical Transactions, 1800, and was read on 26 June, later that year. For electrochemistry, Nicholson and Carlisle's discovery of cathodic hydrogen evolution in electrolysis of water was as monumental as Volta's demonstration of the electrochemical generation of elec tricity described in that letter. The principle was immediately utilized by Sir Humphry Davy who applied it to molten salts produc ing, for the first time, the very reactive alkali metals. From an historical point of view, it seems likely(11) that electrolysis of water was already achieved earlier in an experiment with a large electrostatic generator by Paets van Troostwijk and Deimann in 1789 but separation of hydrogen and oxygen at different poles was not noted. Later, Faraday himself examined the chemical effects of electrostatic electricity and showed them to be the same as those of Galvanic electricity. At that time, it was not possible, from electrostatic machines, to produce a continuous and strong enough current to conduct a significant study of electrochemical reactions as became possible immediately in the 19th century with Volta's pile. As was indicated earlier in this article, a detailed micro scopic understanding of the phenomena underlying electrolysis took a very long time to mature and depended on postulation, proof and acceptance of such concepts as spontaneous ionization in solution, ionic hydration, free ion mobility in solution and the free-electron theory of metals. From a phenomenological point of view, the first quantitative representation of electrolysis as a kinetic process arose from the work of Tafel in 1905 at Erlangen(12) on cathodic H2 evolution. Tafel had been active also in the field of electro-organic chemistry, an area that had been already extensively investigated in earlier years of the 19th century, e.g. by Kolbe, but more from the prep arative than the kinetic-mechanistic directions. Tafel observed that an extra driving voltage, the overpotential η, was required to cause electrolysis to proceed at appreciable net rates expressed in terms of current-density, i. His empirical equation representing this behavior

is fundamental in the kinetics of electrode processes. In Equation 1 a and b are empirical constants: a represents the overpotential at unit current-density (1 A cm-2) while b represents the slope of the increase of overvoltage with logarithmically increasing currentdensity, dη/d log i. The theoretical significance of the logarithmic increase of η with current-density was not understood until much later in the present century in terms of activation ideas in chemical kinetics, starting with Arrhenius in 1889. However, the proper representation of this behavior in electrode kinetics was not formulated until the independent works of Butler (15) in England in 1924 and of Volmer (14) in Germany around 1930 (see later), some 20 years after Tafel's empirical relation for electrode-kinetic behavior.



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Some authors have attributed this delay to the great influence exercised by the electrochemical thermodynamic work of Walther Nernst, commencing in 1889 and continuing onwards for many years. This work emphasized information that can be obtained from measure ments on electrochemical cells at equilibrium, with virtually no reference to the kinetic-mechanistic aspects of the processes invol ved. In fact, it was not until 1924 that Butler (15) gave a kinetic derivation of the Nernst equation, based on the potential-dependence of rates of the backward and forward directions of the process at equilibrium, and 1936 that this approach was dev eloped further by Butler (13) for the reversible H2/H+ elec trode. The components of the energy of the overall process, corres ponding to the Gibbs energy of the electrode reaction that deter mines the Nernst reversible potential, were also treated in some detail in several papers by Butler, e.g. refs. 13 and 15. The Energetics of Electrochemical Charge-transfer Processes It was mentioned in the preceding section that the (standard) Gibbs energy, Δ G ° R , for an overall process of charge transfer at an elec trode, can be considered in terms of component energy contributions. Thus, for a metal ion (M z + ) deposition process

the essential component energies can be visualized through the following Born-Haber cycle of processes that are equivalent to the overall process:

where Gohyd is the Gibbs energy of hydration of the Mz+ ion (a large negative quantity), I M / M z+ is the ionization energy of M atoms, ФM is the electronic work function of the metal M and Golattice is the Gibbs energy for lattice formation of bulk M. ΔGoR can therefore be represented by

Equation 4 applies to the hypothetical single-electrode inter face reaction (2). In an overall, experimentally accessible, cell reaction, it is to be noted that the work functions of the metals involved as the electrode materials cancel out when the metal/metal contact in the external circuit is taken into account. An important development in the 1930's was the representation of the energetic course of electrochemical reactions, especially the step (I) of hydrated-proton discharge in the cathodic hydrogen


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evolution reaction, as a function of the course of the reaction along the "reaction coordinate" from reagents to products. In the hydrogen evolution reaction, the steps are

followed by either


or

the latter process being attributed to Tafel as a rate-controlling step at some metals, e.g. Pt. In steps I, II, or III, the immediate product of discharge is regarded as the H atom chemisorbed at a site on the metal electrode surface, as first recognized by Horiuti and Polanyi in 1935 and treated theoretically by Butler in 1936 (13). The course of the proton-discharge step in electrolytic H2 evolution was first considered in a semi-quantitative way by Gurney and Fowler(16). The interaction of the proton with water, as in H3O+, was represented as a function of distance from the 0 atom by a potential-energy function, L. In the course of neutralization of H 3 O+ by an electron transferred from the metal electrode, the energy Ф has to be supplied (see below) and, at the same time, the proton charge in H3O+ loses its hydration energy, leaving a free H atom in a repulsive relation to the conjugate water molecule left as the pro duct of the discharge process. In Gurney's representation, the possible binding or adsorption of the discharged H atom at the metal was not taken into account (cf. ref. 13) . In electrode kinetics, as empirically represented by Tafels equation, a basic feature is the potential-dependence of the reaction rate (current-density). This effect arises in Gurney's representa tion in a fundamental and general way: as the electric potential V, of the electrode metal is changed by ΔV relative to that of the sol ution (in practice, measured relative to the potential of a reference electrode at open-circuit), the effective value of the electron work function Ф of the metal is changed according to

This has the effect of raising or lowering the energy level of the initial state of the reaction [here H3O+ + e(M)] through change of the energy level of the sub-system of electrons in the metal relative to their hypothetical zero kinetic-energy level in vacuum. That is, the "electron affinity" of the metal can be regarded as being changed by changes of applied potential according to Equation 5. Apart from representing the course of an electrochemical chargetransfer process in terms of a reaction energy-profile diagram, Gurney brought into consideration two other fundamental ideas invol ved in the treatment of electrode kinetics(16): a) introduction of wave-mechanical tunneling of the electrons involved in the process; and



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b) consideration of the distribution of energy levels of elec trons in the metal (i.e. in modern terms, the band structure in relation to the Fermi level from which Ф is measured relative to vacuum) and of unoccupied and occupied orbital states in the reag ent(s), here H3O+. These two factors formed the general basis of all succeeding treatments of electron transfer in electrode processes. Another general condition arises indirectly from (a) in Gurney's treatment, that is, that the electron transfer takes place as a radiationless transition. Thus, the e transfer takes place from an occupied energy level in the metal, at or near the Fermi level, to an unoccupied level in the reagent (usually the lowest unoccupied molecular orbit al, LUMO, in the case of a cathodic process) at the same energy as the level in the metal from which the electron originates. For an anodic process, the opposite considerations apply. The energy-level distribution factor is now recognized as a fundamental factor in the quantum-mechanical representation of elec tron-transfer rates both in heterogeneous redox reactions(7) at electrode surfaces and in homogeneous ones in bulk solution, as well as at semi-conductors(18). One factor of importance for electrode-kinetics was, however, omitted in Gurney's treatment: the role of adsorption of the dis charged H atom at the metal electrode surface. Neglect of this factor led Gurney's treatment to predict activation energies for pro ton discharge that were far too high and independent of the electrode metal. This adsorption energy factor, A, was taken into account in the 1936 paper of Butler(13) who showed how it could account for sub stantial differences of activation energy for proton discharge at Ni compared with Hg. In fact, in this paper, can be traced the birth of the important field of electrocatalysis. The H atom adsorption energy, A, is also a distance-variable in the potential-energy diag ram representation of proton discharge and must, like L, be repres ented as a function of distance from the metal surface by a poten tial-energy function, e.g. a Morse function for the quasi-diatomic molecule MH, representing the adsorbed state of H at M. The role of the H adsorption energy factor in the kinetics of steps I, II and III of the H2 evolution reaction was taken by Parsons(19) in 1958 in a quasi-thermodynamic way, using kinetic equations at the reversible potential. In that paper he showed how the exchange-current density parameter, i o , characterizing the kin etic reversibility of the reaction, depended, in a volcano-like plot, on the standard Gibbs energy of chemisorption, ΔGOH, of the H at the electrode metal through the equilibrium coverage, θH, by H and the corresponding free-site fraction 1-ΘH. Maximum activity arises when ΔGOH = 0. An interesting development from Gurney's paper(16) and from a paper of Bell(20) published soon after, was the application of quan tum-mechanical tunneling concepts to the transfer of the proton in the H2 evolution reaction by a tunneling mechanism, as treated in a paper by Bawn and Ogden(21). Further treatments of this kind were made later by Conway(22), Conway and Sa lomon(23), Christov(24), and Bockris and Matthews(25). However, despite attempts to detect proton transfer by the tunneling mechanism by means of electrochemical ex periments at low temperatures(23) (e.g. down to 180 K ) , no clear


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experimental basis for such behavior could be found. Attempts have also been made to detect such effects in proton-transfer-controlled homogeneous reactions, e.g. processes in which enolization is a rate-controlling step; unfortunately, in most cases, indications of proton-tunneling in the kinetics have been ambiguous.


Kinetic Representation of Electrode Processes Following the seminal paper of Butler (37) in 1924 on the kinetic basis of Nernst equilibrium potentials, an electrochemical rate equa tion was written by Erdey-Gruz and Volmer (14), for a net currentdensity i, in terras of components of i for the forward and backward directions of the process. They recognized that only some fraction (denoted by a or 3) of the electrical energy change ηF associated with change of electrode potential, would exponentially modify the current, giving a potential-dependent rate-equation of the form:

where k and k are electrochemical rate constants for the forward and backward directions of the process, containing exponential terms in the reversible potential, and (3 and a are charge-transfer factors that must be taken as ca. 0.5 in order to represent correctly the experimental value of dr|/d 1n i, as was recognized by Gurney but apparently almost as an afterthought by Fowler (16). Equation 6 refers to a process that is first order in each direction with an electron number n=l. In order that Equation 6 gives rise to the equilibrium potential when zero net current flow takes place (i=0), it is necessary that a = 1-B. Also, in Equation 6, the overpotential η is defined as the difference between the actual electrode potential, E, required to pass a density of current i A cm 2 and the reversible potential, Erev, when i=0 and r|=0; i.e. n = E - E r e v . Equation 6 is referred to as the Butler-Volmer equation. Normally, for significant overpotentials, either one or the other of the two terms is dominant, so that the current-density exponentially increases with η, i.e. 1n i is proportional to BηF/RT in the case, for example, of appreciable positive η values. Here the significance of Tafel's b coefficient (Equation 1) is seen: b = dη/d 1n i = RT/3F for a simple, single-electron charge-transfer process. Equation 6 represents the experimentally realizable possibility, in electrode kinetics, of changing the rate of a reaction over many magnitudes simply by turning the knob of a potential-controller. For more complex processes, e.g. in the case of the desorption steps II or III in the H2 evolution reaction, b can take limiting values of RT/(1+3)F or RT/2F, depending on the mechanism. The evaluation and interpretation of the Tafel b coefficient has played an important role in elucidation of mechanisms of multi-step electro de processes involving coupled charge-transfer and surface-chemical steps, and many cases have been worked out (26). Extension of the Gurney-Butler treatments of the kinetics of electrochemical charge-transfer was made in terms of the "transitionstate" theory of Eyring, Glasstone and Laidler in a paper (27) by


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these authors. This provided the formal basis for representation of electrode reaction rates in terms of the standard Gibbs energy of activation and a transmission coefficient K usually taken, at least until recently, with the value of unity:

where C R is a local reactant concentration at the electrode surface and k, T and h have their usual significance. i is the currentdensity for a z-electron transfer process proceeding at a metal/sol ution potential of ΔФ. Correspondingly Equation 7 can be expressed in terms of exponentials in -ΔHo≠/RT and ΔSo≠/R. The parameter a (or B) in Equation 7 was clearly recognized in the Eyring, Glasstone and Laidler paper(27) as representing the pot ential at the transition state as a fraction of the overall metalsolution potential difference. When the transition state is "sym metrically situated", then a or B is 0.5. There is an interesting relation between the effect of changing electrode potential on the rate of an electrode process and the result of changing acid or base strength (pK) on the rate constant, k, of homogeneous proton-transfer processes, as expressed by the BrФnsted relation: Δln k = a Δln K, where a is BrФnsted's coeffi cient, ca. 0.5. Thus, the value of a or B in the Butler-Volmer equation expresses the fraction (a, B = 0.5) of the applied change of electrical energy, r)F, that is effective in modulating the Gibbs energy of activation of the reaction. The analogy of the electro chemical a or 3 to the BrФnsted coefficient for linear free-energy relations in proton transfer was already recognized by BrФnsted him self with Ross-Kane(32), and in papers by Frumkin(33) and the present author(34). The Tafel relation, especially for electrochem ical proton transfer at Hg, in fact, provides one of the best examples of a linear free-energy relation over a wide range of changes of energy corresponding to ca. 9 decades of change of rate constant. In recent work it has been found(35) that B is dependent on temperature, an effect that can be explained formally in terms of the electrode potential, or interphasial field, affecting not only the energy of activation, through change of the electron's energy, but also the entropy of activation(36) through changes of the inter phasial environment of the reacting particle, e.g. H3O+. This historical account of the development of ideas and experi ments on charge transfer in electrochemistry should not conclude without reference to the Faraday Discuss ion (28) in 1947, held at the University of Manchester. This discussion marked an important turn ing point in electrode kinetics towards more modern and quantitative analyses of electrode process mechanisms and utilization of relative ly new (for that time) techniques, e.g. a.c. impedance studies in the papers by Randles (37) and by Ershler (38). It also brought together many European electrochemists, following the war years, during which little scientific intercourse had taken place on fundamental aspects of electrochemistry.


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Electron Transfer in Inorganic Ionic Redox Reactions Much of the earlier work on charge-transfer processes was directed to studies of the water electrolysis processes which involve "atom transfer" as well as electron transfer, and to some electro-organic reactions. Around the mid-1950's, a more generalized approach was made to the problem of electron transfer, separately by Hush(29) and by Marcus(17,30), who considered the conditions applying to electron transfer between the conjugate ions of a redox couple, both at an electrode and homogeneously. A central concept in this approach(30) was the "reorganization energy", A, associated with change of the solvated state of one ion or the other of the redox couple when electron transfer takes place. The activation energy in this approach is determined, in part, by one-quarter of the quantity X which itself is related to the change of dielectric polarization energy of the ion upon being oxidized or reduced. In its original form, this theory was based on non-specific, long-range dielectric polarization effects, contrary to the types of approach that had been made for the proton-discharge process where short-range, hydration-shell energy changes had been the basis of earlier activation energy calculations(31). The Marcus type theory has had a great influence on the treatment of mechanisms and rates of inorganic inner and outer-shell redox reactions at electrodes and the relations of their kinetics to those of corresponding processes conducted homogeneously. Acknowledgment The author makes grateful acknowledgment to Alcan International and to the Natural Sciences and Engineering Research Council for appointment to the Research Chair in Electrochemistry at the University of Ottawa, and for support of this and related work. Literature Cited 1. 2. 3. 4. 5. 6. 7. 8. 9.

10. 11.

Williams, L. Pearce, Faraday, Michael A biography, Chapman and Hall: London 1965. Thomson, J.J. Phil. Mag. 1897; 44, 293. See Wilson, H.A. Proc. Cambridge Phil. Soc. 1897; 9, 244. Millikan, R.A. e.g. see "The Electron ", 2nd edn., Univ. of Chicago Press: Chicago, 1924. Stoney, G.Johnstone Phil. Mag. 1881, 11, 381. Stoney, G.Johnstone Sci. Trans. Roy. Dublin Soc. 1891, [2] 4, 583. Shuster, A. The Progress of Physics, Cambridge Univ. Press; Cambridge, 1911, p.56. Maxwell, J.Clerk A Treatise on Electricity and Magnetism, Clarendon Press: Oxford, 1873, sec. 259, 260. Conway, B.E. Electricity and Chemistry, Dow Staff Research Lecture, 1966, Univ. of Ottawa, publ. Univ. of Ottawa Press: Ottawa, 1967. Born, M Zeit. Physik, 1920, 1, 45; see also Bernal, J.D. ; Fowler, R.H. J. Chem. Phys., 1933, 1, 515. Leicester, H.M. The Historical Background of Chemistry, Dover Publications: New York, 1956.


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