Con Way 95

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PergamonPqress m Surface Science. Vol. 49, Nu. 4, pp. 331-452, 1995

Copyright 0 1995 Elscvier Science Ltd

OO79-6816(95)00040-2Printed in the USA. All rights reserved.

M)79-6816/95 $29.00

ELECTROCHEMICAL OXIDE FILM FORMATIONAT NOBLE METALS AS A SURFACE-CHEMICAL

PROCESS

B. E. CONWAY

Chemistry Department, University of Ottawa. 10 Marie Curie Street,Ottawa, ON. Canada K 1N 6N5

Abstract

The mechanisms of electrochemical oxide film formation at noble metals are described

and exemplified by the cases of Pt and Au, especially in the light of recentexperimentation by means of cyclic voltammetry, ellipsometry and vacuum surface-science

studies using LEED and AES.

Unlike the mechanisms of base-metal oxidation, e.g., in corrosion processes, anodic

oxide film formation at noble metals proceeds by su&zcr chrmicul processes involving.

initially, sub-monolayer, through monolayer, formation of 2-dimensional OH/O arrays.During such 2-d processes, place-exchange between electrosorbed OH or 0 species on thesurface, and Pt or Au atoms within the surface lattice, takes place leading to a quasi-2-d

compact film which then grows ultimately to a multilayer hydrous oxide film, probably by

continuing injection of ions of the substrate metal and their migration through the growingfilm under the influence of the field.

The initial, sub-monolayer stage of electrosorption of OH involves competitivechemisorption by anions, e.g. HSO,‘, ClO;. Cl-, which inhibits onset of the first stage of

surface oxidation. These processes are demonstrable in experiments on single-crystalsurfaces. The combination of such anion effects with place-exchange during the extensionof the film, leads to a general mechanism of noble metal oxide film formation.

The formation of the oxide films can be examined in detail by recording the

distinguishable stages in the film’s electrochemical reduction in linear-sweep voltammetry

which is sensitive down to OH/O fractional coverages as low as 0.5% and over time-scalesdown to 50~s in experiments on time-evolution and transformation of the states of the

oxide films.

By means of LEED, AES and STM or AFM experiments, the reconstructions and

perturbations (e.g. generation of stepped terraces) which oxide films cause on single-

crystal surfaces can be followed.

Contents

Part 1: GENERAL PHENOMENOLOGY AND EXPERIMENTAL METHODOLOGIES 333

I.1I.2I.3I.4

I.5

I.6I.7

1.x

General Mechanisms of Metal Oxidation and Oxide Film Formation . . . . . . 333Introductory Remarks on Noble Metal Oxide Films . . . . . . . . . 334Early Work on Oxide Film Formation at Noble Metals . . . . . . . . . . . . 335Scope of this Article . . . . . . . . . . . . . . . . 33xRepresentation of Oxide States at Anode Surfaces . . . . . . . . . . . . . . . . . 340

Historical Background of the Development of Experimental Procedures . . . . . 341Theoretical Basis of Electrochemical Techniques . . . . . . . . . . . . . . . . . 343Bases for Charging-Curve and Cyclic Voltammetric Determinations of Surface

331

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332 B.E. Conway

Coverages of Electroactive Species at Electrodes . . . . . . . . . . . . . . . . . .I.9 Electrosorption of Oxygen Species in Successive Stages of Surface Lattice

Occupation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .1.10 Anion Adsorption E ffects in Surface Oxidation of Noble Metals . . . . . . .I. I 1 The Experimental Indications of Anion Adsorption at Noble Metals in

Relation to Surface Oxidation . . . . . . . . . . . . . . . . . . . . . . . . . . . .I. 12 Optical Characterization of Oxide Films . . . . . . . . . . . . . . . . . . . . . . .I.13 Electrochemical Quartz Microbalance Studies . . . . . . . . . . . . . . . . . . . .I.14 The Transition from Reversibly to Irreversibly Electrosorbed OH or 0

Species at Pt and Au . . . . . . . . . . . . . . . . . , . . . . . . . . . . . . . . . . . . . .I. 15 Hysteresis and Kinetic Irreversibility in Oxide Film Formation . . . . . . . . .

350

355

358

359360

361

364

370

Part II: DETAILED UNDERSTANDING OF THE SURFACE OXIDATIONPROCESSES ATPt.............................................. 376

II. 1

II.2II.3

II.4II.5II.6II.7II.8II.9

Microscopic Progression of States of Formation of the Oxide Film at Pt . . . . . 376

Reversibility and Irreversibility of Various Stages of Surface Oxidation . . . . . . 383Kinetic Interpretation of Shapes of i-V Profiles for Formation and Reduction

of Oxide Films on Pt . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 388The Quasi-2-d State of Pt Oxide Film Formation as Revealed in Reduction . . . 391Apparent Limit of Anodic Oxide Film Formation at Pt . . . . . . . . . . . . . . 396Kinetics of the Oxide Film Growth Processes . . . . . . . . . . . . . . . . . . . . . 399Model of the Thin and Thick Oxide Film . . . . . . . . . . . . . . . . . . . . . . . . 400Development of Thick Film Oxide States at Pt . . . . . . . . . . . . . . . . . . . . . . 401Surface Structure Changes at Pt During Oxide Film Formation and Reduction . 405

Part III: COMPARATIVE BEHAVIOR IN SURFACE OXIDE FORMATION ANDREDUCTION AT Au . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 407

III. 1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 407III.2 Time Evolution of the Initial Stages of Surface Oxide Formation at Au . . . . 409III.3 Physical Basis of the Transformation Process . . . . . . . . . . . . . . . . . . . . . 414III.4 Comparative Cyclic Voltammetry Behavior of Principal Index Planes of Au

in Aq. HClO,, and the Role of Anion Adsorption . . . . . . . . . . . . . . 416III.5 Sequential Stages of the Surface Oxidation Reaction at Single-Crystal

Surfaces of Au . . . . . . . . . . . . , . . . . . . . . . . . . . . . . . . . . . . . . . 419III.6 Concentration (Activity) Dependence of the E, Values as a Basis for

Evaluation of Anion Adsorption with Charge Transfer . . . . . . . . . . . . 422III.7 Anion Chemisorption with Charge-Transfer at Au( 111) . . . . . . . . . . . . 431III.8 Role of Water of Hydration of Ions in the Ad-layer at Electrode Surfaces . . . . 432III.9 Structural Changes at Au During Oxidation as Revealed by SERS . . . . . . . . . 434III. 10 Mechanistic Overview of the Elementary Stages of Oxide Film Formation

on Au . . . . . . . . . . . . . . . . . . . .._ . . . . . . . . . . . . . . . . . . . . . . . . . . . . 435111.11 STM Observations of Topographical Changes at Oxidized Au Surfaces . . . . . 438

Part IV: SOME ASPECTS OF THE SURFACE OXIDATION OF Rh . . . . . . . . . . . 444

Acknowledgments . . . . . . . . . . . . I . . . . . . . . . . . . . . . . . . . . 444

References . . . . . . . . . . . . ” . . . . . . . . . . . . . . . . . . I . . . . . . . . . . . . . 445

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ACAESAFM

CPEEELSESCAESDIADEXAFSFTIRLEEDPZCRHERHEEDRTOSEMSERSSIMSSTMUHVUPDXPS

Electrochemical Oxide Film Formation

Abbreviations

alternating currentAuger electron spectroscopyatomic force microscopy

constant potential electrodeelectron energy-loss spectroscopyelectron spectroscopy for chemical analysiselectron-stimulated desorption ion angular distributionextended X-ray absorption fine structureFourier transform infrared (spectroscopy)low-energy electron spectroscopypotential of zero changereference hydrogen electrodereflection high-energy electron diffractionreplacement turn overscanning electron microscopysurface enhanced Raman spectroscopysecondary ion mass spectroscopyscanning tunneling microscopyultra-high vacuumunder-potential depositionX-ray photoelectron spectroscopy

333

Part 1: GENERAL PHENOMENOLOGY AND EXPERIMENTAL METHODOLOGIES

I. 1 General Mechanisms of Metal Oxidation and Oxide Film Formation

The study of oxide-film formation at noble transition metals provides an important transition

case between pure, 2-dimensional surface processes of OH or 0 chemisorption, and formation and

growth of 3-dimensional bulk oxide films or phases. It will be useful, at the outset, to summarise

mechanisms of metal oxidation leading to oxide film formation, as follows:

i) Direct metal dissolution as solvated (S) cations (in water, at low pH or in some non-aqueous

media):

nS + M + M”(nS) + ze (1)

followed by or coupled with hydrolysis of the cation, leading usually to a poorly adherent oxide

or hydroxide film:

zH,O + M”‘(nH,O) + M(OH),J + zH’ + nH,O (2)

Depending on the metal and other conditions, the film product may be the oxide M(0),,2. At

alkaline pH, depending on the acidity of the hydroxide, a soluble complex oxyanion may be

formed, e.g. with Zn as ZnO,*-:

Zn + 20K -+ Zn(OH),L + 2e (34

Zn(OH), + 20K + ZnO,‘. + 2H,O (3b)

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334 EKE. Conway

Depending on the potential and the metal, the precipitated oxide film may be further oxidized to

an higher oxidation state, if it is adherent.

The oxide or oxyanion species that are involved, as in eqns.(2), (3a) or (3b) are usually

represented in the Pourbaix diagram [65] for the metal. The above mechanism is often referred

to as “dissolution and precipitation” and is common in corrosion.

ii) Direct oxide film formation by “nucleation and growth” at the metal surface without

(significant) dissolution, e.g. as with Ni, forming NiO or Ni(OH), with films up to 2-5 nm in

thickness, and with metals that form compact, anodically formed films dependent on the electrode

potential and field across the growing film as with electrolytic generation of films on e.g. Ta, Zr,

Ti, etc. The formation of oxide films by this mechanism, ii, depends amongst other factors, on

the conductivity (or semiconductivity) of the growing film.

iii) Initial formation of submonolayer, extended to a monolayer, of OH and/or 0 on the metal

surface in the form, initially, of a 2-dimensional lattice array. Continuing oxidation and film

growth proceeds by field-dependent interchange between metal atoms in the surface and OH or 0

species in the ad-layer, the so-called “place-exchange” process discussed in more detail later

(Sections 1.9.1.14).

iv) “High-field” growth of usually insulating oxide films through the “Mott-Cabrera” mechanism

[20,21] by injection of metal ions into the oxide film coupled with their migration through the

oxide film to the oxide/solution interface where they are combined with OK from water.

Alternatively, OH or 0. ions formed from water may migrate in the opposite direction to

combine with field-injected metal cations at the inner metal/oxide interface.

v) Combined mechanism: after significant formation of a thin oxide film by mechanism iii,

mechanism iv may take over, leading to thick-film growth, especially if the resulting film is a

poor conductor that can sustain an high electrode field.

Noble metals become electrolytically oxidized by mechanism iii, with v, except under

aggressive corrosion conditions, e.g. in aqua-regia or fused persulfates. However, some trace

parallel oxidation, forming aqua complex ions in solution may accompany oxide film formation.

I.2 Introductory Remarks on Noble Metal Oxide Films

In the case of Pt the interface of that metal with electrolyte solutions has, in many ways. served

as the “model” surface for studies of electrocatalysis, and hydrogen and oxygen chemisorption in

electrochemical surface science, as has Hg in studies of the structure of the double-layer and the

kinetics of electrode processes not involving strong chemisorption or d-orbital interaction [ 1,2,3].

In the case of Pt, interest in its surface electrochemistry was stimulated by the development of

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Electrochemical Oxide Film Formation 335

fuel-cells in the ‘60’s and ‘70’s, principally utilizing this metal in dispersed form as the catalyst

for H,, small organic molecule and hydrocarbon oxidation, as well as, in some cases, for 0,

reduction catalysis. Work on Pt also stimulated related research on the electrochemical behaviour

of other noble metal congeners principally Rh. Pd, Ru and Ir, and the less catalytic metal Au

which is more analogous in its properties (as a solid metal) to Hg.

In earlier times, however, Faraday recognized the special properties of a Pt surface that had

been anodically polarized. Historically, also, Pt and Au were the subjects of some of the earliest

papers in electrochemical surface science,e.g. in the works of Frumkin et al. [4,5], Butler [6,7]

and of Bowden and Rideal [8,9, IO], in parallel with papers on Hg [ 11,12,13].

Apart from their intrinsic interest as surface electrochemical processes, the initial stages of

formation of oxide films on noble metals are also involved in providing the “elements of oxygen”

in the heterogeneous anodic oxidation of small organic molecules, e.g. HCOOH, H.CHO, CH,OH

and other alcohols, as well as sugars, C,H,, etc., to CO, in aqueous media [ 14,151. Additionally,

when thicker oxide films are generated, they form the electrocatalytic surfaces on which other

types of continuous Faradaic reaction take place, e.g. anodic 0, or Cl, evolution, anodic N,

evolution from azide ion, perdisulfate, S20,‘-, formation, the Faraday-Kolbe reaction [54] of

anodic decarboxylation of aliphatic acids, R.COOH, giving CO, + R,, and other organic

adsorption and oxidation processes, e.g. as studied by Tyurin [ 16 and in 171.

Other oxide surfaces of non-noble metals are also of considerable interest in basic and applied

electrochemistry, e.g. NiO.OH and perovskites for 0, evolution from alkaline aqueous media and

in the “Ni-Cd” battery. Co,O, and COO, are of interest as anodic electrocatalyst materials.

Oxides having layer structures, e.g. V,O,,, MnO,, etc. as well as COO,, are also of interest as

hosts for reversible intercalation of Li ions in non-aqueous Li battery embodiments and a variety

of other oxide materials are of practical significance as battery cathode reagents, e.g. PbO,,

AgO/Ag,O, HgO, MnO,. Consideration of these bulk-type oxides is, however, outside the scope

of this present article.

1.3 Early Work on Oxide Film Formation at Noble Metals

The early observations of Laitinen and Enke [29] and of Vetter, that the charge for surface

oxide formation at Pt (up to a given potential, cu. 1.4 V vs RHE) was about twice that for

reduction, is very difficult to account for. particularly as this behavior was observed in two

different laboratories with the experiments having been conducted in a seemingly careful manner.

Somewhat bizarre theoretical ideas were considered [29] in explanation of these surprising result*:

but are no longer tenable, and the anomalous behavior has not subsequently been reproduced.

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336 B.E. Conway

In our own experience, such a result would be difficult to obtain in a cyclic voltammetry

experiment unless some impurity, irreversibly oxidizable over the surface-oxide formation range

of potentials at Pt (0.85 V to 2 1.4 V, RHE), were present at an appreciable concentration.

Alternatively, an unknown accidental asymmetry of sweep-rates in the anodic and cathodic

directions could cause such an effect but a fortuitous factor of two in the sweep-rate for the

anodic over that for the cathodic process would seem to be unlikely.

In the cyclic voltammetry results of Will and Knorr [30] in 1960, there is no evidence for this

effect nor is there from the early but less quantitative results of Russian work [4] on galvanostatic

charging/discharging at Pt electrodes. The significance of the early cyclic voltammetry results

remains a mystery and such behavior has not been encountered since, to the best of our

knowledge, though the presence of impurities [3 1] can introduce disbalances of anodic and

cathodic charges at Pt and other noble metals but only to extents of 5 -10%.

Conclusions from early works on the state of oxygen films at Pt were controversial. Various

investigators [32-391 deduced that at potentials below cu. one volt (RHE) the oxygen was

adsorbed as “PtO” but above this potential as a mixture of PtO and PtO,. At these levels of

coverage of a metal surface by 0 species, it is, however, inappropriate to assign stoichiometric

formulae for the oxide films until they attain “bulk” thicknesses. However, Anson and Lingane

1331 ormed oxide films at Pt for various times and under various conditions, and then chemically

stripped off the films in a solution of 0.2 M aq. HCl + 0.1 M NaCl, subjecting the resulting fluid

to spectrophotometric analysis for “Pt-Cl” complexes. They concluded that oxides in two states

of oxidation were present, PtO and PtO,, in a ratio of 6:l which was not changed upon use of

stronger oxide forming conditions. Later work has shown (see Section IIA.ii) that the thicker

film containing Pt (IV) can be formed independently of the initial “PtO” film which reaches a

limit in its extent of formation.

Except for formation of thicker films, the states of oxide film formation at Pt are better

represented as surface lattice 0:Pt ratios rather than stoichiometric compounds.

From potential relaxation and potential-sweep curves recorded at Pt in 0.2 M HCl + 0.1 M

NaCl, Weininger and Breiter [38] concluded that Anson and Lingane’s results [33] should be

interpreted in terms of a mixed process involving Pt dissolution from the metal rather than from

the oxide, with “PtO” being reduced to Pt.

Various other workers [6,10,40,41] have favored the view that oxygen resides on Pt surfaces as

adsorbed 0 atoms (this was proposed first by Butler and Drever [6]). However, the more recent

work indicates (see Section 1.14) a transition from an initially 2-d adsorbed state to a quasi-3-d

bulk-like state.

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A rather different interpretation of the behavior of oxygen at Pt was extant at one time through

a series of works by Schuldiner and co-workers who introduced the concept of “dermasorbed

oxygen” to represent the state of 0 atoms that had, in some way, entered the skin (rderma) of the

Pt electrode. This idea seems virtually coincident with the concept of “place-exchange” [ 18,191

referred to in more detail elsewhere (see Section 1.14) in this article. The concept of

“dermasorbed oxygen” was useful at the time [42] and it served to emphasize that anodically-

formed oxide films were somehow different from just chemisorbed 0 on the one hand and bulk

oxide on the other. Later, this idea seemed to be translated into a view that 0 atoms actually

dissolved into Pt [39], making a range of solid solutions but later works did not provide evidence

for such a situation. However, Hoare [39] has argued from various literature sources [cf. ref. 431,

related to work function changes at Pt, that 0 may become dissolved in “massive Pt”. In that

work, Kalish and Burstein [43] in fact were able to electrochemically detect 0 supposedly

diffusing out of Pt at which 0 was previously adsorbed. About 3 equivalent layers of 0 could be

thus detected. In Rutherford back-scattering work by Norton et al. [44], sorbed 0 atoms were

detected below the Pt metal surface atoms.

Similar conclusions were made by Schuldiner and Warner [42] in work using rapid and slow

galvanostatic charging and discharging experiments. With rapid charging only the “surface” 0

was measured but, with slow charging, 0 in the Pt surface and dissolved beneath it could be

measured. However, it now seems that such results could be equally well explained in terms of

place-exchanged 2-d films being formed together with a thicker film formed at higher potentials

as discussed in Section 11.4.2. However, it must be stated that Schuldiner’s very careful work on

0 adsorption and reaction at Pt (and his related work on electrosorbed H) provided a solid

pavement for the later, modem work described in other sections of this review.

The suggestion of Vetter and Berndt [45] that the 0 resides on the Pt surface as a peroxide

species has not received general support; most work has favored, rather, a state of bound atomic

0 (or initially OH species), although peroxide intermediates in 0, evolution and reduction have

been postulated. Peroxide species are usually only found under conditions where molecular 0, is

being reduced at non-catalytic metal electrodes although adsorbed peroxy-species have been

suggested as intermediates in the anodic 0, evolution reaction and in persulfate formation.

The spontaneous interaction of noble metal surfaces with molecular 0, has attracted much

interest as this process is closely connected with the potential set up at fuel-cell 0, or air

cathodes. Two questions have arisen [46-481: a) does O2 dissociatively chemisorb at ordinary

temperatures from aqueous solutions and b) what is the significance of the potentials of cu. 0.9 V

(RHE) that are usually set up e.g. at Pt, rather than the thermodynamic potential of 1.23 V (RHE)

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338 B.E. Conway

for the standard conditions. Only at Rh treated in concentrated HNO, has the 1.23 V potential

been realized [48] but Bockris and Huq [49] showed that anodic and cathodic Tafel relations for

the 0, electrode process in very clean solutions at Pt, when extrapolated to zero overpotential,

gave a reversible potential close to the theoretical thermodynamic value for the standard state.

The 0.9 to 0.95 V potential, commonly observed on open-circuit at fuel-cell Pt cathodes, is

probably a mixed potential involving either the O,-reaction process together with oxide-film

reactions, or coupling with the OJperoxide [39] couple. (Electrocatalytic 0, reduction at various

cathode surfaces commonly involves an experimentally detectable peroxide intermediate). The

superoxide ion, O;, can also be an intermediate, probably chemisorbed.

The open-circuit processes nvolving 0, have been examined in a series of papers by Hoare [46-

48]. The dissociative chemisorption process was indicated from results obtained by Bockris et al.

[50] at Pt with various partial pressures of 0,.

When the potential of 0.9 to 0.95 V is set up spontaneously on open-circuit at Pt in a solution

containing 0, near 1 atm. partial pressure, it must be presumed for thermodynamic reasons that

an oxide film is established that is the same as that which would be generated under potentiostatic

anodization at the same potential. From the latter type of experiment, it is known that such a

film would be the 2-d one at substantially less than full coverage by 0 or OH species (O,,+l at

1.1 V, RHE) rather than the phase oxide type that grows only at more positive potentials [see

Sections 11.7,11.8] nd/or during long durations of polarization.

I.4 Scope of this Article

The various sections of this Review provide a comprehensive account of the studies that

have been made of the phenomenology and mechanisms of anodic oxide film formation which

provide an arena for examination of the transition from sub-monolayer, 2-dimensional (2-d)

surface processes involving formation of adatom (“OH” and 0 species) arrays, through quasi-3-d

film formation to development of thicker, bulk-type oxide films that are visible, e.g. at Ru or Lr

under the SEM or optical polarization microscope.

Further, the development of thin oxide films at noble metals involves interesting surface-

chemical processes such as reconstruction and place-exchange [18,19,51], and formation of so-

called hydrous-oxide structures [26], comprehensively reviewed by Burke in the latter reference.

Coupled with the early stages of sub-monolayer array formation involving OH and 0 species,

are competitive processes of anion chemisorption, e.g. of HSO;, ClO,, Cl-, Bi, I-, etc. which

influence the thermodynamics and kinetics of the early, low-coverage stages of anodic oxide film

formation, as will be described later.

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Another important aspect is that the continuous Faradaic processes of anodic Cl,, Br,, N, and 0,

evolution go on always (except in some special absolutely anhydrous solutions in solvents such as

CF,COOH [52] or CH,CN (531) on an oxidized surface of the noble metal, even at Au where the

onset of oxide film formation is not until some 0.12 V positive to the OJH,O reversible potential

(E” = 1.23 V). Hence, the electrocatalytic surfaces for these anodic reactions are the external

interfaces of the oxide films with the electrolyte solution. This is an important practical and

fundamental factor in considering anodic electrochemical reactions that take place at noble metals.

The scope of this article can be defined, in one way, in terms of discussion of the factors that

determine the initial stages of electrochemical oxide film formation at metals as surface

processes:

(a) role of solvent, usually water, adsorption at the metal surface;

(b) anion chemisorption from the electrolyte solution;

(c) 2-dimensional (2-d) electrodeposition of arrays of OH and 0 species on the metal surface

amongst an overlayer of adsorbed solvent and anions:

(d) reconstruction or place-exchange [ 18,191 between the OH or 0 species and surface metal

atoms, initially in 2-d arrays, into a quasi-3-d surface phase or domains of such a phase,

accompanied by desorption of the previously chemisorbed anions, and

(e) growth or thickening of the oxide film by the “high-field” types of mechanism proposed by

Mott and Cabrem [20] or by Ghez [21], as referred to earlier in Section I.l(iv).

At base metals in alkaline solution, steps (c) and (d) are usually by-passed by nucleation and

growth [22] of 3-d oxide phases, initially as islands on the metal surface, but the initial nucleation

event, often a stochastic one, will commonly be a localized “surface electrochemical” process.

This is the case for metals at which oxide/hydroxide film formation does not take place by the

“dissolution-precipitation” type of mechanism (see Section 1.1).

ln outlining the scope of this article, it will be useful to indicate what areas will not be treated,

at least in detail: these cover (a) electrocatalysis and electrosynthesis at oxidized noble-metal

anodes; (b) thick-film oxide formation at noble metals, especially Ru and Ir and (c) industrial

processes utilizing noble metal or oxidized noble metal anodes. Also, we shall try to avoid

repetition of material already reviewed in the excellent earlier articles by Woods [23] and by

Dignam [24] on metal oxide film formation, and in the general review on electrochemical surface

science by the present author [25]. More recently. the topic of thick, hydrous oxide films at

noble metals has been reviewed in detail by Burke [26].

Finally, we can summarize the scope of this paper as covering the following topics:

a) the surface chemistry of the initial stages of oxide film formation at noble metals, especially

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340 B.E. Conway

Pt and Au, including the role of co-adsorption of anions and place-exchange;

b) the transition to thick film formation and its independence from formation of the initial 2-d

film;

c) the kinetics of growth of thicker oxide films;

d) the mechanisms of oxide formation and reduction, and the origin of irreversibility between

these processes;

e) the quasi-2-d behavior of thick oxide films at Ru and Ir; and

f) the optical reflectivity behavior of oxide films.

In order to present a structured view of the oxide-film formation processes that arise at Pt or

Au, it will be most clear (if not most logical) if we anticipate, in some ways, and take into

account, general conclusions that can be reached from an overview of the literature on Pt and Au

oxide film formation and reduction, especially the more recent works [55,56,57]; on this basis, it

seems that anodic thin-film oxide growth on these metals must be treated in terms of four

distinguishable stages:

i) Anion adsorption that arises prior to onset of OH/O electrodeposition (equivalent to

competitive “pre-oxidation” of the metal surface by anion chemisorption with charge-transfer

IhO,hll);

ii) Electrodeposition of OH or 0 species in a sub-monolayer (array), competitively, amongst the

pre-adsorbed anions [56,57];

iii) Conversion of this low-coverage 2-d OH or 0 sub-monolayer to a quasi-2-d film by place-

exchange [ 19,51,58,59], formed up to a limit of cu. 880 $ cm-* at Pt and a lower limit at Au.

This process can be accompanied by anion desorption (ClO,-, HSO,. but not with Cl- or Br-);

iv) More or less independent [62] continuing growth of a porous, hydrous oxide layer on top of

the quasi-2-d film in the case of Pt and in a related, but not identical, way at Au (e.g. a visible,

yellow, thick film of oxide can eventually be formed at Au under strongly anodizing conditions

[63.64], and it can even be seen to peel off!).

It will be seen, however, later, that while the general behavior of Au is similar to that of Pt and

Rh (and also to the initial stages of oxide film formation at Ir and Ru), the details of the behavior

at Au are significantly different, possibly because lattice displacements of atoms in Au are more

facile than with Pt, Rh, Ru or Ir which are harder metals having larger lattice energies and the

strength of anion pre-adsorption is also different.

I.5 Representation of Oxide States at Anode Surfaces

In so far as this review is concerned mainly with 2-d aspects of surface oxide film formation, it

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is necessary to comment on the representation of oxide-film species as stoichiometric compounds

such as “AuO”. Au,O,” at e.g. Au anodes. as is to be found in the earlier literature [63,68] and

some more recently [69]. Because such formulations can be misleading, we suggest that they

should be avoided except when oxide films are thick enough to behave as true and

thermodynamically well defined stoichiometric phases; otherwise, a metal:oxygen surface ratio

for the 2-d species should rather be specified, as exemplified by the notation of Parsons and

Visscher [30] for the 2-d oxide state that is formed to a limit corresponding to “PtOPtO” or 880

pC (real cm)-’ in reduction (oxidation state of Pt, +2). For sub-monolayer array states, a lattice

occupancy ratio should be specified relative to the single-crystal plane on which the electrosorbed

species is deposited, e.g. as in refs. [55] and [57] on our own work.

Thermodynamic considerations would suggest that a characteristic new oxide “phase” in 3-

dimensions on a metal surface should hardly be expected unless at least two layers of “M,O,”

structure are sandwiched between further surface layers of such species which would

have Gibbs energies different from those for the outer “sandwich layers”. Thermodynamically,

a 3-d bulk phase is recognized when its bulk Gibbs energy per mole is appreciably less than its

surface Gibbs energy. The latter is usually dependent on coverage. In such cases of “bulk”

films, the oxide reduction should be associated with at least cu. 1800 pC cm-* (depending on the

stoichiometric numbers x and y). Also only for such true 3-dimensional phases, is a single-

valued potential for reduction to be expected as in a galvanostatic experiment, or a narrow spike

as in a linear-sweep voltammetry experiment. Most surface oxide species generated initially as

very thin films on noble metals, are not of the latter 3-dimensional category.

I.6 Historical Background of the Development of Experimental Procedures

Following the use of the “charging-curve” method for measurement of the double-layer

capacitance at Hg by Bowden and Rideal [80], this procedure was independently applied to the

study of electrochemical processes at Pt and Au by Frumkin and co-workers [4] in the USSR and

by Butler and co-workers [6,7,12,4 I ,8 ] in England in the early ‘30’s. These works were the first

to distinguish. at Pt. the processes of electrochemical H adsorption and ionization near the H,

reversible potential and the reactions of surface oxide formation at Pt [4,7,83] and Au [41,83],

and reduction both at substantially more positive potentials (>0.9 V,RHE) than those for H

ionization. In the work of Pearson and Butler 171,a cathode ray oscilloscope was employed

(probably for the first time in this field of work) to record rapid transients generated at high

charging currents to minimize unwanted effects due to back-diffusion of H, in the electrolyte to

the electrode.

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342 B.E. Conway

The use of the charging-curve method was continued into later times in the high quality work uf

Schuldiner and Hoare and co-workers in various papers, mainly at Pt and Pd. In this work. the

concept that the oxide film at Pt was a “dermasorbed” layer was enunciated [42]. This was a

useful idea and anticipated the more specific model of the oxide film in terms of place-exchange

[ 1X] following the initial development of partial coverage by 2-dimensionally chemisorbed OH

and/or 0 species 1551.

In the work of Armstrong. Himsworth and Butler [X3]. the charge passed over the potential

range C’U. 0.95 to 1.35 V, RHE, was for the first time interpreted as due to formation of a

monolayer of 0 atoms on the (polycrystalline) Pt surface.

Information on electrochemical surface processes at Pt and other noble metals made a majot

step forward through the use of potentiostats in the work of Will and Knorr in 1960 [30] where

the potential-sweep method. already in use in polarography at Hg (Sevcik [75]), was applied to

the study o f such interfacial processes at Pt and other metals. This work opened a new vista of

opportunities for research in electrochemical surface science especially as the method gives a

d~fi~rrnriol response to the effect o f modulated potential through the recording of currents t=CQs

associated with charging of the adsorption pseudocapacitance C, in parallel (Section 1.7) with the

double-layer capacitance. Thus the sequence of surface processes associated with “charging” of a

Pt electrode interface could be much better resolved than in the procedure of “charging-curve”

transients which is essentially an irztrgrui method. Kozlowska and Conway [X4], however,

overcame the latter limitation by developin g a direct analogue procedure for rl#erenriari~~~ of

(constant current) charging curves by means of an operational amplifier with an output to an

oscilloscope.

Although a potentiostatic voltage control system was already in use in the ‘30’s by Hodgkin and

Huxley [X5] in their work on squid axon nerve “action potentials” and Hickling [X6] had

described a potentiostat in 1940, it was the advent of the German Wenking potentiostat [cf. ref.

301 capable of voltage control to 1 mV in a response time-scale of a few us, that enabled major

advances tu be made after 1960. Such systems could be used in a programmed way by coupling

with electronic multiple function generators. Later developments, into the present period, allow

the same operations to be made under precise computer control, with digital data acquisition and

data processing, especially using fast digital oscilloscopes for primary data recording.

The use of controlled-potential techniques that are available through electrochemical

instrumentation. provides an important degree of control in studies of (anodic) oxide film

formation at metals that is not possible in examination of oxide film formation at metals from the

gas phase nor so easily by means of controlled current methods. In the former case. however.

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variations of 0, gas pressure, p, over a wide range would be analogous to variation of electrode

potential, E, (AE P -A(RT/zF) In p). In both cases, variation of temperature has major effects on

the kinetics of oxide film formation and reduction processes.

1.7 Theoretical Basis of Electrochemical Techniques

Anticipating that this article will be read not only by electrochemical specialists but also by

some “vacuum” surface scientists, it will be useful to give a brief account of the electrochemical

procedures involved in electrochemical surface-science studies since much of the material which

follows is based on results from such experimentation. In a general sense it is to be noted first

that electrochemistry is basically a “%-dimensional” science although quasi-3-dimensional

electrode systems can be achieved with porous or powder electrodes, as in fuel-cell and battery

technology.

Other general and particular aspects of electrochemical surface science were reviewed in some

detail in an earlier article by the present author in this journal [25].

In electrochemical processes, an additional variable of major importance is the electrode

potential. E, which can be employed to modulate the state of the system either thermodynamically

or kinetically. From a thermodynamic point of view, the electrochemical potential, p, of a

charged species, i, is defined through the relation

h = p,’ + RT In ai + ze$ (4)

where a, is the activity of species i in the given phase and C$s the inner potential of the phase,

relative to @= 0 at an infinite distance from the phase. ze is the charge borne by the particle so

that ze$ is the electrical contribution to the Gibbs energy of the species i.

Application of relation (4) to a charge transfer process at equilibrium, e.g.

S + M * MsZ+ + ze (in M) (5)

where S is a solvent, leads to the Nemst equation for the hypothetical single’ electrode potential

of M in equilibrium with its cation Mz+ in solution in S. viz.

E, = E,” + (RT/zF) In ai (s) (6)

where E”, like uiO, s some standard-state value, usually defined for a, = 1. In eqn. (6), E, is

determined by the difference between &, and & where M and S refer to the metal and solution

phases, respectively.

When the electrochemical equilibrium involves 2-d chemisorption of an electrochemically-

*Only the difference of rwo single electrode potentials, together with a contact potential in the externalcan be experimentally measured in electrochemical experiments (see ref. [70] for further discussion).

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344 B.E. Conway

generated chemisorbed species, e.g. H produced in a process such as

M + H,O+ + e * M.H,, + H,O (7)

l-0, C,+ E 0,

application of equations like (4) and (6) lead to

8,/l-8, = KC,+ exp [-EF/RT] (8)

in the case of a B-independent equilibrium constant, K, for the adsorption process (7). Eqn. (8) in

an electrochemical Langmuir isotherm for electrochemical adsorption of H by charge transfer in

reaction 7.

Expressed logarithmically, eqn. (8) takes the form of a Nernst equation for the surface process

(7), viz.

E = E, - (RT/F) In Cn+ + (RT/F) ln [8,/l-8,] (9’)

where E, is related to K.

For conditions where K depends itself on 8, e.g. due to lateral interaction effects, eqn. (8) takes,

for the form of Frumkin’s electrochemical adsorption isotherm:

[8$1-$1 exp go = ho Cn+ exp [-EF/RT] (10)

where g represents the pairwise interaction energy in units of RT of the chemisorbed species in a

random array.

If the surface is heterogeneous, so that the (apparent) standard Gibbs energy of chemisorption

(AG” = -RT In K) depends on site occupancy, other types of isotherm apply, e.g. the Freundlich

or Temkin isotherms. Limitingly, the latter form of adsorption isotherm represents 8 as a linear

function of E. It will be seen that relation (10) also limitingly can have the same form - for

intermediate values of e and for sufficiently large g. However, the physical origins of these

limiting cases are quite different but are often confused.

Eqns. (8) or (9) apply to a “single state” of adsorption as 8 varies from 0 -+l. However, it is

often found, especially in electrochemical adsorption of adatoms (so-called underpotential

deposition phenomena [25]), that several stages of the chemisorption can be distinguished below

monolayer coverage, each having a characteristic standard p0 or corresponding E” and behaving in

an approximately Langmuirian adsorption manner. This situation is often even more clearly

discernible at well prepared single-crystal surfaces, [90,94] so is not an artifact due to surface

heterogeneity, though this question was controversial for some years (cf. ref. [71]).

The key aspect of electrochemical procedures is their sensitivity; in charging curve or cyclic

voltammetry experiments (see below), charge can be “counted” in p coulombs cm-’ very

sensitively (-210 pC = 1 monolayer) enabling submonolayer quantities of electrosorbed species to

be determined down to coverages of OS-l% of a monolayer and resolution of energies of

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345lectrochemical Oxide Film Formation

adsorbed states to cu. 2 kJ mole (= 20 mV), or better.

In the case of kinetic aspects of electrode processes nvolving charge transfer, it is the

electrochemical uctivution energy, AC?, that is dependent on potential, viz. (cf. ref. 70):

AG’ = AG*,, f p EF/RT (11)

where p, for simple processes, s an electrochemical Brqnsted-type factor having a value cu. 0.5,

i.e. changes of Gibbs energies of activation for electrochemical charge-transfer processes due to

variation of electrode potential are cu. half the changes of the electrical energy of the charged

reactant particle, normally the electron. Changes of electrical energy AEF, correspond normally

to controllable changes of the Fermi-level energy of the metal electrons relative to “vacuum”.

Since B is 0.5 for a symmetrical energy barrier defining AG’, p is also often referred to as the

“barrier symmetry factor”.

Because AG’ varies often in a linear manner with E for atom transfer processes (but sometimes

in a quadratic manner for simple “outer-sphere” electron transfer reactions) and rates of processes

are exponential in AG’ (Arrhenius equation), it is usually found that the log of rate constants for

charge-transfer processes varies linearly with electrode potential or overpotential, ?l (the

difference between the actual E required for some net rate to be established and the reversible

value of E for equilibrium in the same process). This linear-log relation, originally discovered

empirically, is often referred to as the Tafel equation [72,70].

Correspondingly, electrode reaction rates can be modulated expmzentiully with potential

according to the relation

i(q) = i(q=O) exp [pqF/RT] (12)

where n is the overpotential, E-E,,,, and i(q=O) is the exchange current-density (usually written i,)

corresponding to q=O, and measures the current (density) passing reversibly in each direction of

the electrode process at equilibrium.

A linear form of eqn. (12) arises for sufficiently small 7 values (PqF<RT) when the back-

reaction rate of the process becomes comparable with the forward rate. Then the kinetics are

quasi-Ohmic (i proportional to q) while for Pq/RT > 1 the kinetics are exponential in q (Tafel

relation):

q(i)=a + b In i (13)

=a + (RT/PF) In i (14)

The linear limiting form of the kinetic relation is useful in alternating-voltage modulation (“A.C.”

impedance studies) as the frequency-response of the electrode process can be measured and

analysed in terms of the kinetic relaxation characteristics of the processes involved (so-called

“impedance spectroscopy”).

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346 B.E. Conway

Thus, from the above summary, it is seen how in electrode processes, the electrical variable

gives an additional “handle” for the study of surface processes nvolving charge-transfer, that is

not normally available at the metal-vacuum or metal-low pressure gas interface. However,

because of the necessity for a bulk-phase electrolyte to be in contact with the investigated

electrode surface, many of the elegant procedures of the gallimaufry of high-vacuum surface

science techniques, e.g. Auger and ESCA surface analysis, in situ LEED and RHEED

characterization of surface structures, EELS, ESDIAD, SIMS, etc are inapplicable. The recently

developed procedures of direct observation of actual lattice geometries by STM and AFM are,

however, applicable through a solution.

Despite these limitations, employment of “transfer” techniques in the use of UHV instruments

has, in recent years, enabled virtual in situ examination of electrode surfaces by procedures first

developed by Hubbard. This has added a valuable arsenal of techniques for application in

electrochemical surface science.

Some optical procedures can, however, be applied directly in situ with electrolyte present, e.g.

single reflection FTIR spectroscopy [73], surface-enhanced Raman spectroscopy [74], specular

reflectivity, second-harmonic reflection spectroscopy and ellipsometry. Other modem techniques

such as EXAFS and related techniques, and direct surface X-ray diffraction are also applicable.

The most spectacular in situ techniques, recently applied to the examination of electrode surfaces,

are of course, scanning tunnelling microscopy (STM) and atomic force microscopy (AFM), as

mentioned above.

The most sensitive direct electrochemical technique for examination of both the thermodynamic

and kinetic aspects of electrochemical surface processes s the method of linear sweep

voltammetry or, in a cyclic embodiment, “cyclic voltammetry”. The methodology was derived

from that in polarography, especially fast-sweep polarography of SevCik [75], as mentioned

earlier.

In this technique, which is especially useful for non-continuous, 2-d Faradaic electrochemical

surface processes such as 2-d atom array formation and atom desorption (by ionization

electrochemically), the potential is modulated over a controlled range of values by means of a

potentiostat (operating as a potentiodyne) addressed by a linear-sweep generator, linearly in time

in one direction of potential change or another, in reverse. In fact, the potential addressed to the

electrode (relative to the potential of some appropriate reference electrode in the system) can be

programmed in various more complex ways to provide a “conditioning program” operating on the

electrode.

The current response to this modulation of electrode potential is recorded conjugate to the

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varied potential, giving rise to a so-called voltammogram or cyclic-voltammogram (see Fig. 1).

2 x d.t. charging

3current

t

I

Potential, E>

Fig. 1 Typical cyclic voltammogram for an electrochemical surface process (response currentdensity, i, plotted as a function of potential, E, cycled between two limits at a rates = dE/dt).

Resolution of the initial stages of oxide film formation and reduction, and of multiple stages of

adatom deposition below monolayer level, have depended very much on the application of cyclic

voltammetry and, in a related way, on potential-modulated reflectivity measurements, including

ellipsometry. We hence explain briefly the basis of the cyclic voltammetry method.

The current response of an electrochemical surface process to a linear variation of potential

sweep can be represented (cf. refs. [76-781) as follows: both the forward and backward directions

of the process, e.g. as for reaction (7), should be considered. The rate of the forward process

depends on the free-site fraction, l-8, and the rate of the back reaction on the occupied-site

fraction. 8, as well as on the time-variant potential, E(t) = East, where s is the rate of modulation

of potential in V s-‘:

Forward current density: 2 kz (1-e) exp [P(E+st)F/RT] (15)4-t

Backward current density: i = k 8 exp[-( I-P)(E-cst)F/RT] (16)

where c is the concentration of the source of the adspecies involving a reagent in the electrolyte

solution, e.g. H,O+ (for H), solvated metal ions for M atom deposition, etc.

In considering eqns. (15) and (16), it should be noted that surface processes cannot proceed

continuously in time, except when multi-layer film growth can take place, as is actually the case

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348 B.E. Conway

with oxide formation; that is, the rate of (16) becomes 0 as 8 + 1 and the rate of (16) also

becomes 0 for 8 -+ 0, so that a maximum current arises at 9 = 0.5 (in the reversible case, see

below).

Only when the sweep-rate is small enough in relation to the k values does the electrochemical

surface process remain almost at equilibrium [77,79] throughout the sweep, i.e. both the forward

and backward rates in eqn. (15) and (16) determine the net rate of deposition of the ad-species,

i.e. i = i - i. When the sweep-rate exceeds a critical value, s,, which characterizes the limit of

reversibility of the process (the process remaining kinetically reversible for s<s, or sas,),

polarization sets in and current-peak potentials shift with log s as in a Tafel relation [79] (eqn.

13).

When SX,, either the forward or the backward direction of the process (depending on the

diractinn of the sweep) is dominant so that i + i or i + i alone; then the process is kinetically

irreversible and the dependence of 8 on E(t) is determined either by eqn. (15) or by eqn. (16)

[76,77]. The transition between the reversible and irreversible behavior is shown in Fig. 2

according to the calculations of Srinivasan and Gileadi [76].

Fig. 2 Theoretical current response curves in cyclic voltammetry for a single-state surface processshowing transition from reversible to irreversible behavior with increase of sweep-rate(KS, to SX,). (From Srinivasan and Gileadi, ref. 76).

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In the kinetically irreversible regime, the peak potentials, E,, move proportionally to log s with

a Tafel slope RT@F or RT/(l-P)F, depending on the direction of sweep. Under equilibrium

conditions (s<s,), eqns. (15) and (16) correspond, for i = i, to the electrochemical adsorption

isotherm (eqn. 8) and to the “reversible” curve of Fig. 2. Then each value of E, in the course of

its change with time at a raters,, is thermodynamically related to a corresponding value of 8

according to an isotherm of the type (8) or (10).

Under reversible conditions, an adsorption pseudocapacitance arises, characterizing the

reversible surface process. We designate it C+,, and it corresponds to the derivative of 0 with

respect to E, i.e. the derivative of the electrochemical adsorption isotherm, e.g. eqn. (8) or (10).

In the case of (8),

c&I= s, t&,/c-)=gI F/RT Xc exp [EF/RT]

(1 + FC exp [EF/RTJ2 1where 0, is the coverage generated under reversible conditions (i = i) and q, is the charge

required to generate full coverage by the ad-species, e.g. 210 uC cm-* for H in reaction (7) at a

Pt( 111) surface: for a surface with N substrate atoms per cm’

q, = NF/N, cm-* (19)

where N, is Avogadro’s number and F = 96,500 C mole-‘.

Eqn. (18) obviously .gives a maximum in C,, and it arises at a characteristic peak potential, E,,

corresponding to 8 = 0.5 and detemined by K. Alternatively.

C,,r = [q, F/R-U@,(1-W (29)

which has a maximum value q,F/4 RT, ca. 2200 p cm-‘. Eqn. (18) is also the relation for the

differential electrochemical isotherm for process (7). As discussed in Section 1.8, C, is easily and

directly determined at any potential from the relation i = (C,) x (s) in a voltammogram.

For irreversible conditions (cf. Fig. 2), a pseudocapacitance, C,, still arises since 8 depends still

on E but according only to either eqn. (15) or eqn. (16). The Er, s then not at 8=0.5 and the

maximum in C, is also lower (see Fig. 2).

Generally, current-profile (C,) peak widths (e.g. as measured by the width in potential, AE,, at

the half-height) indicate the extent and type (repulsive or attractive) of lateral interactions in the

film: a broad peak indicates repulsive interactions, a narrow one, attractive effects (771. A very

sharp peak usually indicates a phase-transition, or formation or reduction of a 3-d phase having,

limitingly, a singular - valued Gibbs energy.The above relations (18) or (20) form the basis of application of cyclic voltammetry to

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350 B.E. Conway

electrochemical surface processes except that, in practice, the approach to full coverage with

changing E usually takes place in two or more distinguishable stages corresponding to different

states of array formation as the available surface progressively becomes occupied by the ad-

species. Order/disorder transitions or 2-d phase transitions are believed to occur. In the case of

oxide formation processes, the current response to increasing potential corresponds initially to

monolayer formation, then to place-exchange and film-thickening at higher positive potentials.

I.8 Bases for Charging-Curve and Cyclic Voltammetric Determinations of Surface

Coverages of Electroactive Species at Electrodes

In order to assist non-electrochemists who are concerned with surface science in following the

significance and operational principles of methodologies of electrochemistry as applied to

electrochemical surface science, it will be useful to outline the bases of the procedures that are

involved in processing results derived from use of the electrochemical methods used in this field:

they are essentially simple and quite sensitive.

The basis for charging-curve evaluation of coverages by electro-chemisorbed species is simply

the application of Faraday’s laws on a microscopic scale. For an arrest in a charging curve (e.g.

Fig. 3) lasting, say, for 2.1 s at a constant current-density of 100 pA cm-‘, the charge q passed for

deposition or desorption of an electroactive species such as H, or OH is = q = ji.dt = 210 uC cm-*.

It is easy KO alculate that this corresponds to ca. 1.30 x 10” atoms cm.‘, e.g. on a Pt(ll1)

surface, i.e. a monolayer.

Potential / V, E,,

Fig. 3 Typical “charging curve” for deposition of surface H and OH species at Pt, together withcharging of the double-layer capacitance. (Based on ref. 7).

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In the case of cyclic voltammetry, the integral of the i(V) profile over the voltage range of a

current peak gives the charge passed during passage of potential through the range V,-V, say,

where the current response i(V) arises. Thus

CT=

s'i(V) .dt =

"I s"2l(V!/S.dV i21)"I

where s is the voltage sweep-rate, s = dV/dt. Since i = dqtdt it is seen that the quantity i(V)/s is

a capacitance. It corresponds, in a useful way, to the differential coe$icient of the

electrochemical adsorption isotherm for the electrosorbed species since the coverage fraction. 8,

of such a species arising in a Faradaic surface reaction is related to the charge passed q(V) for its

electrodeposition or electrodesorption of the species according to

W) = qWYq, (22)

where q, is the charge for formation of a monolayer (f3 = l), eqn. (19).

The above relations apply to the determination of coverages by electroactive species provided

that:

a) corrections are made for charge passed in charging of the double-layer:

(23)

over the potential range V,-V, under examination in the charging transient,

b) no (or negligible) Faradaic charge for continuous simultaneous processes, e.g. H, or 0,

evolution, passes. This condition is normally met by conducting the cyclic voltammetry

experiment between +O.O5V and ca. +1.4 V, RHE (thermodynamically, to 1.23V at 298K) in the

case of regular aqueous solutions containing no other species reducible or oxidizable over that

range.

In the case of oxide film formation, C,, cannot usually be assumed to be the same as for an

oxide-free surface of the metal. C,,, for such conditions, may be obtained in favorable cases by

means of impedance measurements at high frequency in a suitable potential range, where no

electroactivity of the oxygen species of the oxide film itself arises in response to fast a.c.

modulation. Normally this must be done at potentials well displaced from those at which the

very initial stages of oxide film formation (see Section 11.2) arise, since at some metals e.g. Pt,

Au, such processes can be kinetically reversible.

It is to be noted that cyclic voltammetry i(V) profiles represent the differential of the chargingcurves since the current generated in a sweep is the response function to dV/dt and involves the

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352 B.E. Conway

total capacitance C (double-layer plus that arising from potential-dependence of coverage by

electrosorbed species), i.e.

i(V) = dq/dt 2 C(dV/dt) (24)

and C is a differential quantity involving coverage:

c = c,, + c, (25)

= C,, + q,(dWdV) (with dq = q,d@. (26)

Cyclic voltammetry i(V) profiles (cyclic voltammograms) provide more detailed information on

surface electrochemical processes than do the corresponding “integral” charging curves.

However, the latter can be differentiated, during recording, by means of an operational amplifier,

as demonstrated by Kozlowska and Conway [84] as mentioned earlier; the resulting curves then

give equivalent information to those from cyclic-voltammetry although the latter method probably

gives more precise differential information. integral charging curves often obscure important

details of electrosorption processes, as a function of coverage.

Thus, for example, direct charging curves for oxide film formation on Pt or Au usually have

given the impression [41,81,83] of a straight line, the slope of which is apparently a constant

(reciprocal) capacitance. Cyclic voltammograms, on the other hand, clearly show that the

capacitance (eqn. 24) is not constant and important details of the potential dependence and hence

that of 8 are revealed [30,55,77].

The sensitivity of electrochemical procedures for examination of chemisorption processes with

charge transfer is easily demonstrated by a simple calculation involving application of Faraday’s

laws on a microscopic scale to coverage changes at electrodes. Thus, for adatoms deposited with

passage of one e per atom, e.g. H on Pt, and taking approximately 1.7 x 1Or5atoms per cm2 on a

Pt( 100) surface, the charge, q, required (cf. eqn. 19) is

q = q, 5 1.30 x 10” x 96,500/6 x 10” C cm-’ = 210 uC cm-’ (27)

This is a quantity that is easily determined with a 1% accuracy, or better, by electrochemical

measurements, e.g. cyclic voltammetry, involving integrating current with respect to time, q =

]i(t).dt over the potential range, AE (or V,-V,), covered at a sweep-rate s V s-‘, i.e. over a time

interval AV/s in the integral.

The potential sweep method is specially convenient for studying the sequential stages in

development of surface oxides on noble metals as the potential is increased in the anodic

direction over successive cycles [55,56,57]. It allows. like the opening of a fan, the separation of

the partial processes of an overall complex reaction on a time-dependent potential scale according

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to the rate constants of the various processes in relation to the sweep-rate, s, used. By choice of

the proper s, not only can the partial processes be resolved but some slow reactions can be

eliminated or their products observed in a controlled way. Therefore the dependence of the i vs

V profile on the sweep-rate can be diagnostic [77] of the reactions occurring on a surface.

However, it has to be kept in mind that all resulting information is pertinent specifically to the

time-scale of the sweep-rate used.

In this method, each energetically different electrochemical reaction on a surface is registered on

a potential (and time) scale in the form of a current peak. For surface reactions, under

equilibrium conditions, a specific isotherm will represent each particular energy of adsorption or

deposition of species on the surface and will correspond to an individual peak in the i vs E

profile. From the differences of the peak potentials, AE,, that arise when the surface processes

are reversible, (KS,,), the differences of the free energies, AE = -A(AG)zF for the particular

reactions are obtained.

For deposition of the same species, e.g. OH or 0, from a given solution, the differences of the

energies usually arise from different energetic states of the surface sites, related to different

emerging d orbitals. This may be due to different symmetries of arrangement of atoms on the

surface coupled with resultant lateral interaction energies, or may be created by formation of

distinguishable sublattices by the deposited species themselves or by overlay lattices formed by

species already adsorbed on the surface, e.g. anions.

For irreversible processes, kinetic information can be obtained from the change of E, with s

[76.77,79]. The i vs E profile can supply further information about the surface reaction such as

the charge, q, passed during the reaction from the “area” under the peak, about interactions

between the species on the surface from the shape and “half-width” of a peak [76,87], and finally

the pseudocapacitance, C,, connected with the surface reaction [87] or, in the case of simple

charging of the double-layer, its capacitance, C,,, can be obtained from the ratio i/s at high s

values.

When the valency of the reacting species is known and full charge transfer occurs, the number

of molecules reacting and the number of occupied sites of each particular energy type can be

easily evaluated from the respective charges under the individual peaks. The situation is more

complicated when the extent of charge transfer [60,61], Se, is not known a priori. Then, for the

full stoichiometry of the reaction, knowledge of 6 or the number of molecules per site is required

from other data.

The main difficulty in obtaining all the required information about a particular surface reaction

from i vs V profiles along lies, as for the other “spectroscopic” methods, in the resolution of

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354 B.E. Conway

overlapping peaks since the differences of energy between various types of s ites are not large, so

that the resulting peaks usually overlap. Sometimes they can be “kinetically” separated by

variation of the sweep-rate [76,77] or application of superimposed fast modulation [SS].

We conclude this section with a summary of the types of information obtainable with cyclic

voltammetry and related techniques (Table 1).

I.

2.

3.

4.

5.

6.

7.

X.

9.

10.

11.

Table 1. Sensitivity and Applications of the Cyclic Voltammetry Technique for

Characterization of Electrochemical Surface Processes

Sub-monolayer surface array states can be distinguished with standard Gibbs energies

differing by cu. 50 mV (~4.5 kJ mol.‘) or less.

Standard Gibbs energies of surface array states can be evaluated with an accuracy of 20

mV (~1.8 kJ mol.‘), depending on resolution.

Differential adsorption isotherms for the ad-species can be directly measured in terms of

the potential-dependence of the pseudocapacitance, C,, determined from i/s.

Coverages or changes of coverage in sub-monolayer ad-states can be evaluated to an

accuracy in 8 of 0.5 to l%, or 5%. depending on resolution.

Coverages of ad-species down to 0.5% in 8 can be evaluated.

Coverage changes in time can be followed down to 20-50 us in pulse transients.

Longer term changes of coverage, i.e. by oxide species can be accurately followed.

Reversible surface processes can be clearly distinguished from irreversible ones and the

rate-constants for the latter can be quantitatively determined: also the condition for

transition from reversible to irreversible conditions in the kinetics can be evaluated.

Optical reflectivity measurements can be coupled with linear sweep voltammetry, enabling

the optical properties of the varying sub-monolayer films to be correlated with charges for

film formation and reduction or desorption, and with effects due to competitive anion

adsorption and irreversibility.

Response currents in cyclic voltammetry and potential-pulse experiments

(chronoamperometry) can be easily measured at the sensitivity level of 1 uA and potentials

can be measured and controlled to 1 mV. Integrated charges can be evaluated with an

accuracy of 2 uC or 1% relative accuracy, corresponding to 0.5- 1.0% in 0.

Reversible electrosorption processes can be distinguished by their response to fast

modulation [88] both in linear sweep voltammetry and in relative reflectivity experiments,

as well as in A.C. impedance spectroscopy.

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1.9 Electrosorption of Oxygen Species in Successive Stages of Surface Lattice Occupation

Early studies of surface oxidation of Pt by the constant-current charging method [41,83] seemed

to indicate a relation of constant slope between potential and quantity of charge passed, i.e., a

behavior corresponding to constant adsorption pseudocapacitance (C,) associated with

electrodeposition of OH and 0 species up to and beyond a monolayer of 0 species (cu. 1.40 V,

RHE). However, under linear potential sweep conditions, it is clear (see Fig. 4 and ref. 55) that

C, is not constant with increasing potential beyond that corresponding to significant

commencement of OH deposition.

(Ill1 100) “ROH”800 IV to"1200 --

AEp=50mV AEp=75mV AEp=lOOmV

Fig. 4 Distinguishable, but overlapping, current peaks OAl, OA2, OA3, in cyclic voltammeuy atPt in clean aq. H,SO, at 298K (from ref. 55).

Working in very pure dilute solutions of HClO, and H,SO, in pyrodistilled [31] water,

Kozlowska, Conway and Sharp [55] showed that three distinguishable but overlapping current (or

C,) peaks were observable (Fig. 4), as can also be found at low index surfaces of single-crystal

Pt.

In this work, these authors (551 suggested, for the first time, that concepts of surface chemistryshould be used to understand this behavior, i.e. that the distinguishable but overlapping peaks

corresponded to successive stages of 2-dimensional occupancy of the Pt surface lattice by OH or

0 overlay arrays which they wrote for a (100) surface as PtJOH, PtJOH and Pt/OH (see Fig. 5).

Corresponding possibilities for occupancy by 0 species at higher potentials as OH species

become converted to O’s with loss of a proton and electron, can be envisaged.

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356 B.E. Conway

LATTICE

P1,OH

Fig. 5 Succession of lattice arrays of electrosorbed OH or 0 species on (100) or (Ill) substratesingle-crystal surfaces (from ref. 55).

The above designations, PtJOH etc., do not represent stoichiometric compounds but rather the

lattice occupancy ratio by OH to Pt atoms. On the (111) surface, ratios PtJOH to Pt/OH would

arise.

These stages of 2-dimensional surface lattice occupancy find parallels in the multiple state

chemisorption of H on single-crystal surfaces of Pt [89] and of Pb adatoms on single-crystal

principal index faces of Au, studied by Schmidt et al. [90], where the successive stages of surface

lattice occupancy are much better resolved into clear peaks at characteristic electrode potentials,

i.e. Gibbs energies of electrosorption.

While the OH and 0 electrosorption peaks for Pt are always found to overlap considerably, at

Au single-crystal surfaces, studied by Kozlowska, Hamelin and Conway [56,57], several clearly

distinguishable peaks arise below monolayer OH coverage, and their forms and separations on the

electrode potential scale depend substantially on the geometry of the substrate surface lattice, e.g.

for (1 I 1 , (110) and (100) surfaces (see Part III). A minimum in C, is usually characteristically

observed 1911 at the level of monolayer (Au/O) coverage by 0 species just prior to

commencement of anodic 0, evolution; in fact this minimum has been suggested by Burshtein et

al. [91] as a basis for real surface area determination of Au electrodes (the “Burshtein minimum”

method) since the formation charge for the 0 monolayer on (111) Au is 400 pC cm’ to this

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minimum, corresponding to one 0 per Au surface atom. (At Au, real surface areas cannot be

determined by accommodation for adsorbed H as they can at Pt [92]).

It must be mentioned that the surface occupancy ratios, shown above, are only nominal values

for successive array structures since it is now well known from vacuum surface science

experiments [93] (LEED and STM) that 2-dimensional array structures are often non-epitaxial

(so-called “non-commensurate”) with respect to the substrate surface lattice and may have a

“rotational” relation of their geometries to that of the underlying metal surface lattice. Such

behavior is commonly also associated with “reconstruction” processes in chemisorption.

Further, it should be emphasized that the multiple state chemisorption phenomena, often

observed in electrochemisorption processes n underpotential adatom array formation, is not

primarily due to polycrystallinity of the surface [cf. refs. 71,901, though this may complicate the

observation of multi-state, sub-monolayer chemisorption. In earlier works on electrochemical

adatom deposition it was commonly thought that polycrystallinity was the reason for observation

of multiple-state chemisorption of simple monatomic adsorbates, e.g. H on Pt, Pb on Au.

However, the careful work on well characterized single-crystal surfaces, carried out later by

various authors [e.g. 89,90,94] has shown clearly that resolution of multiple states, distinguishable

by cyclic-voltammetry peaks at various potentials in cyclic-voltammograms, was usually b~trfr

seen at single-crystal surfaces than on polycrystalline ones of the same metal. The arguments in

support of this position (which is now experimentally established) were presented critically and

comparatively in several papers in the 70’s by the present author with Kozlowska et al. [71]. Of

course, even the most carefully polished and annealed single-crystal electrode surfaces have some

remaining surface concentration of step edges or kink sites and other defects in their “2-

dimensional” structures, as is now directly visible under scanning tunnelling microscopy.

However, it is clear that such small concentrations of surface irregularities as can now be directly

seen at atomic resolution under the STM, could not account for the major distinguishable features

that are observed in cyclic-voltammograms [cf. refs. 56.57,89,90] of adatom deposition and

desorption processes on various noble metal substrates.

From such results, it can be concluded by inference, but with little doubt, that the multiple

states of submonolayer oxide film formation that are quite generally observed at the noble metals

arise for similar reasons to those for metal adatom deposition, though it must be recognized that,

in the case of OH and 0 electrosorption, reconstruction processes set in already at low (cu. 5 to

1.5%)coverage and complicate the situation, leading (see Sections I.14 and 11.2) o irreversibility

between the processes of anodic electrodeposition (of OH and 0) and cathodic reduction of the

oxide films. It is important to stress that this is a general phenomenon of some fundamental

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358 B.E. Conway

significance and is also observed in the interaction of low-pressure 0, gas with metal surfaces

f951.

1.10 Anion Adsorption Effects in Surface Oxidation of Noble Metals

The onset potential of surface oxidation of noble metals is very sensitive to the concentration

and type of anions of the electrolyte in which the surface oxidation studies of noble metals are

conducted. Halide effects were first demonstrated by Breiter [96] but quantitatively examined by

Bagotzky et al. [97] and by Conway and Novak [98] at Pt. Kozlowska, Conway and Hamelin

[56,57] suggested a general and fundamental basis for such effects including those of ClO; and

HSO; ions, as will be described in Section III.3 for Au.

First it is necessary to outline the phenomenology of anion chemisorption effects at electrodes.

This was examined and treated by Grabame [3] for the case of the Hg electrode. The hierarchy

of adsorption effects of anions at Hg is illustrated in Fig. 6 in terms of the shift, AEpzc,of the

potential of zero charge (P.z.c.) of Hg by various anions (an effect also dependent on their

chemical potential in solution, related to RT log a, their activity, a. The adsorption of anions at

Hg involves their polarizability and Gibbs energy of hydration, properties related to their ionic

radii and their electron-pair donicities [99], the latter being related to their electronic and, for

polyatomic anions, their geometric structure. The same factors determine their adsorbability at

noble metals. Since their chemisorption is related also to the electron affinity of the electrode

metal, it is dependent on the metal and its electrode potential relative to the p.z.c., and thus on

the electronic work function. Since anion adsorption also requires, in some degree, displacement

of previously adsorbed and oriented [ 1001 solvent dipoles, this introduces a further metal-specific

factor.

Bockris et al. [ 10 ] considered some of these factors quantitatively and made estimates of

energies of adsorption of ions. However, since the net, relatively small energy of adsorption is

determined by the sums and differences of larger quantities, a reliable evaluation is quite difficult,

especially in regard to partial dehydration effects and partial charge transfer [61] related to

electronic interaction with the metal surface through the ion’s donicity. A difficult problem is

how to deal with the image charge interaction that an ion experiences at short distances from

metal surfaces (102, 103; reviewed in 104). The Gibbs energy of ion adsorption is a more

difficult quantity to evaluate since it requires a calculation of the entropy of the system.

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Fig. 6 Hierarchy of increasing strengths of chemisorption of anions at Hg in terms of shifts ofpotential of zero charge (Ap.z.c.) in 0.1M solutions.

I. 11 The Experimental Indications of Anion Adsorption at Noble Metals in Relation

to Surface Oxidation

Phenomenologically, specificity of anion adsorption at the noble metals is clearly discernible

from the effects anions have on: a) the potential for onset of surface oxidation currents, e.g. in

linear anodic sweep experiments [105]; b) changes in the current profile for the surface oxidation

as a function of potential; c) related to (a), the reversibility of deposition and reductive

desorption of OH at low coverages (see Section II.2 and Fig. 13); d) the form of the cathodic

current profile in reduction of the oxide film (dependent on the extent to which it has been

previously formed under anodic polarization) and e) the optical properties of Pt and Au surfaces

(Section 1.12).

Several examples of the specificity of anion adsorption effects at polycrystalline Au and Pt are

shown in Fig. 7 and Fig. 14 in relation to the oxide film formation and reduction current profiles.

Anion effects are also observed on the H deposition and desorption profiles at those metals (Pt,

Rh, Pd, Ru, lr) where underpotential deposition of H arises (not at Au). Both types of effect

arise on account of competition between chemisorption of the anions and that of OH or H

species.

Kozlowska, Conway and Hamelin [56,57] have pointed out (see Section 111.3) hat the effects of

anions on the initial stages of 2-dimensional oxide film formation are to be treated in terms of a

preoxidation of the metal surfaces since the anions that have competitive effects on the initial

stages of OH electrosorption, e.g. at Pt or Au, are chemisorbed with an appreciable degree of

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360 B.E. Conway

charge transfer already at potentials significantly below those at which electrosotption of OH

from the H,O solvent commences. Such chemisorbed anions therefore have an electrosorption

valency [60] number that is an appreciable fraction of their formal charge.

Log (Concn. / mol dm-3

)

Fig. 7 Specificity of anion chemisorption effects at polycrystalline Pt in relation to onset ofsurface oxide formation processes (from ref. 184).

In alkaline solutions, of course, formation of the oxide film takes place directly from OK ions

of the electrolyte so there can be no competition from other anions of the solution. However, at

Au in alkaline electrolyte there is a clear indication [ 1051 of a reversible pre-adsorption of OH-

with significant charge transfer over a potential range of cu. 0.5 V. Beyond this potential range,

regular 2 or quasi-3-dimensional surface oxide film formation takes place with characteristic

irreversibility. At Pt in alkaline solution, such behavior is not observed but the first 15% of OH

coverage is, however, more or less reversibly deposited [55] but only over a potential range of cu.

0.15 v.

I. 12 Optical Characterization of Oxide Films

Optical methods, such as infra-red absorption, surface-enhanced Raman spectroscopy and

reflectivity techniques involving direct and modulated reflectance, ellipsometry, second-harmonic

reflectance, provide ideal, in situ procedures for characterizing electrode surfaces and species

chemisorbed thereon since there is normally no perturbation of the surface except at high photon

energies where U.V. photo-electron excitations could arise, depending on the band-structure of

oxide films. With oxide films at noble-metal electrodes, methods based on changes of relative

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reflectivity and of the ellipsometry parameters, A and W, have been mainly used. They provide

useful information complementary to that derived by purely electrochemical procedures.

i) Principles involved. The principal optical method for examination of properties of oxide

films at metals and semiconductors is ellipsometry. In particular, ellipsometry and relative

reflectivity measurements provide one of the few non-electrochemical methods that can be applied

at electrodes, in situ, in the presence of the electrolyte solution. Parameters of interest that

characterize the properties of the film are a) its average thickness, d; b) its refractive index, n

and c) its optical absorption coefficient K. d is related to the electrochemically measurable charge

for film formation or reduction, and n and K are related to the electronic properties of the film

(5 1,70, 12.113,l 14). The equations relating A and w to d, n and K, and the general

phenomenology of ellipsometry are to be found e.g. in the monograph of Wolf [106] or in the

review articles of Gottesfeld [107] and of the present author [ 1081, and in the early work of

Tronstad [ 1091.

Modern instrumentation, developed by Ghan with Rudolph Instrument Corp., N.J., allows A

and w measurements to be conducted dynamically on a time-scale down to cu. 50 ms, so that

optical studies on kinetics of growth and reduction of oxide film can be followed. Also,

measurements can be made at vtious wave-lengths including the infra-red range (“ellipsometric

spectroscopy”) [ 1 o] using the Rudolph S 2000 “automatic” ellipsometer.

The ellipsometric parameters, A and w, are derived from a Fourier transform of the digitized

periodic intensity I according to the relations:

tan%= (a#) + 1 (28)

cos A= -a& tan w (29)

where K= (b, + 2a, -$)/2 (30)

1.13 Electrochemical Quartz Microbalance Studies

The extraordinary mass sensitivity of the piezoelectric, quartz-crystal micro-balance resonator

(ECQM) to material adsorbed on its surface has been known for some time [123]. Applications

in electrochemistry to electrosorption from solution, with the quartz crystal immersed in solution,

have been demonstrated [ 124- 1291 and have provided a new tool for in situ studies of

metal/solution or electrode/solution interfaces. For example, single UPD atomic layer formation

of Pb and Ag on Au, has been studied [ 128,129].

The procedure has been applied to the examination of oxide films on Au [ 125,126] and an

illustrative record of the mass response is shown in Fig. 11 from Stiickel et al. [ 1271.

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362 B.E. Conway

Fig. 8 Relative intensities of p (0) and (s) (+) components, and A(A) as a function of potential forlight reflected twice from a Pt electrode in 0.5 M aq. H,SO, (from ref. 119).

i d)

k

Fig. 9 Complex refractive index (n = n-i@ for the oxide film on Pt in 0.5M H,SO,. Shadedareas show domains of values compatible with the optical data for four wave-lengths: a)x=250 nm; b) 209 nm; c) 350 nm and d) 400 nm. (Points x indicate assumed refractive

index of Pt). (From ref. 119).

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Pt. Aq. H,SO,

REDUCTION FROM. +,48” A

+1.22v 0

+,.,ov .

+,.02v +

I Cathodic Peak

-------einodic-polng~line

eoxta [e / Pt atom]

Fig. 10 Plot of changes 6 of A in reflectance measurements at Pt in the oxide formationregion as a function of potential (RHE) or corresponding Q,,K& ratio (from ref.120)

(b

Fig. 11 Mass responses at a Pt electrode due to oxide formation, measured by means of theECQM (cf. Stiickel et al., ref. 127. From M. Zhang, Ph.D. Thesis, University of

Ottawa, 1994).

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364 BE. Conway

The mass change is determined through changes, Am, of the resonator frequency; however,

although quite significant and easily recordable changes of frequency arise on oxide formation,

the interpretation of results is rendered difficult on account of the following matters:

a) incorporation of water into the oxide structure, especially with hydrous oxide structures;

b) potential-dependent anion adsorption or incorporation and c) binding of water on the exterior

surface of the oxide film. and its interaction with adjacent bulk water. Usually the mass change

scales with the charge for formation or reduction of the oxide film at Pt so that the instrument

gives an integral response like a charging curve. However, electronic or digital differentiation of

the frequency change signal could be made with respect to potential. Like the fact that the hound

did not bark, in the “Hound of the Baskervilles” detective story, perhaps the most significant

result at Au is that no net weight change accompanies oxide film formation at that metal, a

puzzling but interesting result found in the work of S tiickel et al. [ 1271, contrary to the behavior

found at Pt. This effect may be due to simultaneous desorption of preadsorbed HSO; as OH/O

species are competitively deposited.

At the moment, it seems that the ECQM experiments do not provide any essentially new results

or details on the process associated with oxide film formation that are not already known through

the application of cyclic voltammetry or potential and wave-length scanning ellipsometric

procedures. Perhaps by much more detailed comparisons between weight change, and formation

and reduction charge evaluations as a function of potential, this method will yield some

interesting and original results. Certainly, the incorporation of water and anions in growing oxide

films. or initial desorption of anions in the initial stages of film formation, is a matter of great

interest but quantitatively meaningful results on these matters are still lacking. Why there should

be qualitative differences between the scaling of (frequency change) Am with weight changes at

Pt vis ~1vis Au is quite unclear but could be associated with the known strong adsorption of

HSO, at Au which becomes desorbed upon oxide formation at that metal, giving fortuitously,

perhaps, virtually zero mass change upon oxide film formation, as noted above.

1.14 The Transition from Reversibly to Irreversibly Electrosorbed OH or 0 Species

at Pt and Au

That some degree of reversibility arises between oxide formation and reduction at low extents of

oxide development was already noted by Gilman [ 1301. This is manifested as a

pseudocapacitance in a.c. modulation [ 1321 at low potentials (0.8-0.95 V), which is appreciably

larger at low frequencies indicating a relatively small rate constant for the “reversible” process.

As noted in Section 1.12, Barrett and Parsons [ 13 ] made modulated reflectivity measurements

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366 B.E. Conway

0.5 M H,s04, S=3.5 V S-’

4

2

0

2

4

Potentia I E/V

Fig. 12 Resolution of a “kinetically reversible” response in the initial stages of oxide filmformation at Pt to fast modulation, superimposed on a slow linear potential sweep

(from ref. 88).

Fig. 13 Resolution of an “envelope” of reversible anodickathodic current responses fromirreversible behavior in cyclic voltammetry at Pt in strong aq. HClO, at -60°C

(from ref. 56).

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Under such conditions of low temperature, the place-exchange rearrangement process is slowed

down thereby better revealing the initial adsorption process. This is clearly seen in a sequence of

cyclic voltammograms at Pt taken to successively more positive potentials in sweeps, each

followed by a corresponding cathodic reduction sweep taken at a rapid rate to minimize

opportunity for place-exchange to occur. The resulting series of cathodic curves displays an

envelope of current response for the irreversible reduction of surface oxygen species clearly

distinguishable from the one for reversible reduction of the species electrosorbed at less positive

potentials (Fig. 13).

In aqueous alkaline solution (Na,CO, or NaOH), some indication of separability of the

reversibly from the irreversibly reducible initial oxide film is already discernible at room

temperature. This behavior can be understood in terms of the diminished tendency for

competitive anion adsorption to arise in alkaline than in acid solution since the potential of zero

charge of the metal is then ca. 0.8 V nearer to the potential for onset of surface oxidation, so

there is less electrostatic tendency for anion adsorption to arise. In OH- solutions, a further effect

arises due to the substantially smaller tendency for OH- (cf. F ion) ion to be chemisorbed

compared with oxyanions such as HSO;, HCO; or ClO.... The adsorption of the HSO; or SO;’

ions has been extensively studied at Pt mainly by the radio-labelling technique of Kazarinov [62]

and more recently to Horanyi et al. [ 1381 and Wieckowski et al. [136]; it was earlier indicated in

a paper by Schuldiner [ 1371.

By more detailed studies at Au electrodes, the role of competitive chemisorption of anions in

relation to the initial stages of surface oxide formation can be clearly demonstrated [ 1051. Fig. 14

shows the comparative effects of several anions, including OK (in Ba(OH), solution, to avoid

carbonate effects due to CO, contamination) on the cyclic voltammetry profiles for surface oxide

formation and reduction at Au at a successive series of increasing positive potentials. The

strongest effects, not only on the oxide formation profiles but also on the cathodic ones, arise

with the borate anion, B,O,*., which is a double-chain condensed polymeric anion. The fact that

the reduction profiles are also affected in a major way may indicate that the strongly adsorbed

anions are incorporated within the oxide film and are not entirely desorbed as the oxide film is

progressively developed.

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368 B.E. Conway

Au/O.1 M HCIO,

Fig. 14 Effects of competitive chemisorption of anions on oxide film formation andreduction at Au in various electrolyte solutions (from ref. 10.5).

The competitive effect of anions on the initial stages of surface oxide formation at Au can be

clearly demonstrated in another way, as shown by Kozlowska et al. [ 1391. Oxide film formation

was again followed by means of cyclic voltammetry but in a supporting electrolyte of 0.1 M aq.

HClO, containing 1O-3M HSO; ion. In a steady sweep from +0.05 V (RHE) through the oxide

formation region, the curves of Fig. 15a were obtained. However, if the potential was initially

held at +0.05 V for a short time (i.e. cu. 0.4 V negative to the P.Z.C.) to establish a surface more

or less free from chemisorbed HSO;, and then shot at 500 V s-’ to the oxide formation region

(thus giving insuff-icient time for anion chemisorption to be established), the curves of Fig. 15b

were observed. It is quite clear that the latter conditions (referred to as “anions off”) allow the

initial, reversible region of electrosorption of OH or 0 species to be well resolved from the

irreversible region at more positive potentials. In the “anions on” condition, corresponding to Fig.

15a, the development of the initial reversible region is hardly resolvable, which is the situation

also for Pt in aq. H,SO, or HClO, at room temperature. The behavior shown in Figs. 15a and

15b is very important as it shows that the initial stage of OH electrosorption must take place

normally amongst a lattice of “pre-adsorbed” anions, prior to their desorption at higher potentials

as OH and 0 coverage is driven to increase.

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i” ’ ,

Fig. 15 Effect of HSO; adsorption in aq. HClO, on resolution of reversibly formed andreduced oxide films at Au electrodes: a) HSO; “on” the surface; b) HSO,“anions off’ the Au surface due to insufficiency of time for their adsorption in afast sweep (500 V s-l) from 0.05 V (RHE). (From ref. 139.)

These results are very important for developing a microscopic model of the sequence of

processes nvolved in the initial stages of anodic oxide film formation at Pt and Au, and probably

also at the other noble metals (see Sections II.1 and III.2 and Scheme 2 in 111.0).

The distinction between a reversibly formed state of the oxide film (in its initial formation, e.g.

at Pt) and that irreversibly formed led [56] to a useful picture of the changing stage of the oxide

film as increased extents of oxide formation takes place. The latter can arise either from

increased positive potential and/or by allowing the film to develop in time [56,115] at a given

fixed potential within the range where oxide film formation is possible.

Reverst ble OH Rearranged PIOH

(al (b)

Fig. i6 Schematic diagrams of increasing extent of deposition of OH or 0 species on a Ptor Au electrode with progressive transformation of an initially 2-d electrosorbedstate to a quasi-3-d, place-exchanged state.

Fig. 16 shows such a picture [56] in terms of the progressive transformation of an initially 2-

dimensional electrosorbed state (like that of H on Pt in the potential range 0.0 to 0.35 V. RHE) to

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370 B.E. Conway

a quasi-3-d state arising on account of place-exchange between the initially electrosorbed OH or

0 species and underlying surface Pt atoms. From the experimental behavior observed it seems

that this transformation is almost totally irreversible. even when formation and reduction of the

film are conducted at very low sweep-rates. However, the behavior is partly kinetic in origin

since slow development of the film proceeds logarithmically in time (see Section 11.6)and once

the film has been formed, its reduction at an appropriate constant potential also proceeds

logarithmically in time. This irreversible transformation is intimately connected with the

mechanism of film growth as a function both of time and potential of the electrode.

1.15 Hysteresis and Kinetic Irreversibility in Oxide Film Formation

The formation of anodic oxide films on the noble metals, beyond the level of the initial sub-

monolayer, is always found to be irreversible. This irreversibility has interesting features: the

greater the extent of formation of the oxide or the higher the potential at which it has been

formed, the l[>wer, less positive, must the potential be in order for it to become reduced. No

matter how slowly is the potentiodynamic sweep conducted in a cyclic voltammetry experiment,

this irreversible behavior is still observed. This type of irreversibility is also clearly discernible in

the “charging curves” recorded in the early works of Butler and of Frumkin referred to earlier.

The behavior described above is contrary to thermodynamic expectations, e.g. for the formation

and reduction of a species successively in two or three oxidation states where the higher oxidation

state is normally reduced at an higher potential than that for the species in the lower oxidation

state(s).

The formation and reduction of oxide films on the noble metals is much more analogous to the

phenomenon of hysteresis, e.g. in magnetization. This analogy is illustrated in Fig. 17 where we

show the charge - potential relations for formation and reduction of anodic oxide films at Pt

generated at successively increasing potentials in relation to the well known hysteresis in

magnetization. Of course, the analogy is only a formal one since the two types of processes

involved are physically quite different. However, hysteresis in adsorption phenomena, e.g. of

gases or vapors sorbable in charcoal, is quite well known [ 140,141] and the isotherms for

desorption (relating quantities sorbed to the gas pressure) and adsorption are often far from

superimposable.

At the oxide films, states of formation generated at the highest potentials are those which are

reduced only at the lowest potentials; thus the phenomenon is closely analogous to what arises in

magnetization where domains which become polarized only at highest field are depolarized only

at correspondingly high fields in the opposite direction. A model illustrating this type of

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situation, also applying to adsorption, was treated by Everett [ 140,141].

The explanation of such phenomena in terms of totally irreversible processes, where the

pathway for the forward (adsorption) process is different from that for the backward process, has

been given by Everett in the above papers in relation to his model referred to in refs. 140 and

141.

In the case of oxide film formation and reduction at noble metals it seems that the process of

place-exchange provides the difference of pathway in formation from that in reduction of the

oxide film, i.e. the reduction is not microscopically the reverse of the oxidation, that is required to

characterize an hysteresis. The co-adsorption of anions also provides some coupled effects in this

difference.

250 550

Chorge /pC cm’

Fig. 17 Increasing hysteresis between the potential ranges for formation of oxide films on

Pt or Au, and their reduction, as greater extents of oxide film formation aregenerated at higher positive potentials. Analogy to hysteresis inmagnetization/demagnetization is to be noted.

We illustrate the hysteresis schematically in Scheme 1 below:

Further oxidation

Scheme 1

For simplicity, anion adsorption and desorption effects (see Sections 1.10 and 111.3) are not

shown in the above Scheme.

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372 B.E. Conway

Of course. no physical or physico-chemical process can proceed (at finite temperatures) at a

zero rate, so, given sufficient time, there can be conversion of the oxide film from its formed

state to reduced state (the metal surface) at potentials in a sweep where reduction seems otherwise

not possible: that is, there are kinetic effects. Thus. we have mentioned in Section II.6 that

oxide reduction at potentials well positive to the current peak in the cathodic voltammogram can

be realized but only at a very slow rate, in fact logarithmically in time, as for the kinetics of

formation of the oxide film beyond the sub-monolayer state (Section I.9), where the growth

process is also linearly logarithmic in time [ 142,143] (Section 11.6).

When oxide films are formed at Pt at elevated temperatures, 100°C or greater in water, the

hysteresis is very much diminished and the potential ranges for oxide formation and reduction

overlap much more as illustrated in Fig. 18. Eventually at sufficiently high temperature, it might

be presumed that irreversibility and hysteresis are virtually eliminated. However, as found for

oxide or other film formation and reduction processes at base-metals, there is always some

overpotential required in the formation reaction and some, in the opposite direction around the

reversible potential, for the reduction process, as illustrated in Fig. 19 for formation and reduction

of PbCl, films on Pb in a cyclic voltammetry experiment.

It is interesting, in this context, that recent experiments in this laboratory on thin-film oxide

formation at Ni and Fe down to low temperature (-90°C in methanol/water mixtures [ 1441) where

the oxide film formation can be limited to sub-monolayers, show that irreversibility is still

maintained, i.e. no reversible region can be observed like that found at Pt at low temperatures

[%I]. This suggests that, at such base metals, the oxide is not formed by an initial 2-d

electrosorption process but rather by a “nucleation and growth” [22,24,145J type of mechanism

which is usually irreversible.

The hysteresis between formation and reduction of oxide films on the noble metals is a

fundamental phenomenological aspect of their behavior and of the mechanisms of the processes

involved.

Two apparent exceptions to the hysteresis behavior we have referred to, e.g. for Pt and Au, arise

in surface oxidation of Ru and Ir; here, however, thicker film formation can arise but the

oxidation and reduction processes, beyond several monolayers, are remarkably reversible.

However, it is now known that the behavior is quite different from that of Pt or Au, since the

oxidation and reduction processes involve not formation of the film and its reduction back to

mrtul but rather redox processes between two or more oxidation states of the noble-metal ion e.g.

RI.?+,RI?, Ru4+ (or even RU~), within an intact oxide film. This result was demonstrated by

Gottesfeld by means of ellipsometry which showed that on reduction of IrO, films, a lower oxide,

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Fig. 1X Improvement in reversibility between oxide formation and reduction processes at Ptin 0.5 M aq. H,SO, at an elevated temperature.

POTENTIAL/V, RHE

Fig. 19 Comparative example of irreversibility in formation and reduction of PbCl, at Pbwhere a bulk phase (PbCI,), thick fiim, is formed. Curves shown for successivesweeps to increasing anodic potentials.

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374 B.E. Conway

rather than the metal surface resulted.

However, initially upon cycling Ru or Ir electrodes, cyclic voltammograms, like that for Pt, do

urisr, i.e. with irreversibility, but after only a few anodic/cathodic cycles to 1.4 V or more,

rrversibk oxidation/reduction processes become observable (Fig. 20) and the oxide films

progressively thicken to levels corresponding to 200-300 times the initial monolayer

formation/reduction charge, the so-called “charge-enhancement factor” [ 1461. Then the oxide

films are visible under an optical microscope with interesting structure under the SEM (Fig. 21).

Because of the facility [ 1471 with which the initial, irreversibly formed, monolayers of oxide at

Ru or Ir become converted to the “reversible redox” type of film, care must be taken in

observations of the initial film formation process. The fact that such initial films can be observed

[ 1471 at Ru and Ir, means that, basically, these metals behave with respect to 2-d oxide film

formation, exactly like Pt and Au which confirms the generality of the mechanisms of the initial

stages of surface oxide film formation at noble metals that we emphasize in this article.

However, at Pt and Au, rather special conditions, e.g. modulation, are required for the

development of thick, hydrous oxide films [26,180].

It seems that the thick “redox-reversible” types of films formed at Ru and Ir must arise from

injection of Ru and Ir ions through the initial 2-d film to build up the “redox-reversible” film that

is not easily reducible to the metal (in fact, at IrO,, the redox process is electrochromic and the

oxide film in its lower state of oxidation is much less conducting than that in the higher state

(ho,)). This process is. in some way, analogous to that at Pt where also thick films of oxide can

be generated on top of the initially formed 2-d film [48].

However, in this case of Pt, such films show no reversible redox behavior and, in fact, their

formation and reduction is associated with the largest degrees of irreversibility or hysteresis.

It is interesting that Ni and Co, which normally exhibit irreversibility in formation and reduction

and their oxide films (at Ni, to NiO.OH or NiO,, at Co to Co,O,), a “redox-reversible” state of

the oxides can be generated by high potential conditioning and then these materia ls, like RuO, or

IrO,, exhibit a redox pseudocapacitance corresponding to “mirror-image”, anodiclcathodic cyclic

voltammograms that are observed at Ru [149] or Ir [ 115,146].

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I

.

I I I I I I I0.2 0.6 1.0 I-4

Potential/V, ERHE

Fig. 20 “Reversible” cyclic voltammogram (due to film redox processes) for a thick RuO,film developed at a Ru electrode upon cycling over the range 0.05 V to 1.40 V

PW .

Fig. 21 SEM photographs of the thick oxide film (several microns) develope

electrode upon extended anodiclcathodic cycling in aq. H2S04.

d at a Ru

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376 B.E. Conway

Part II: DETAILED UNDERSTANDING OF THE SURFACE OXIDATION PROCESSES

AT Pt

11. Microscopic Progression of States of Formation of the Oxide Film at Pt

Starting in the early ‘70’s, a number of papers appeared [55-591 that substantially extended the

understanding of the details of the mechanism of surface oxide fiim formation and reduction at

Pt. including aspects of the interesting irreversibility that arises between these two processes

beyond an extent of formation of electrosorbed OH species corresponding to 10 to 20% coverage,

depending on pH and electrolyte anion concentration.

In two important papers by Vetter and Schultze [58,59], potentiostatic and galvanostatic pulse

measurements were employed to characterize the progression of the states of formation of the

surface oxide film with increasing potential and the related behavior in reduction. These authors

distinguished the initial formation of a chemisorbed film of (supposedly) 0 species from

subsequent, more extended formation of a “regular”, place-exchanged “O-PC* layer followed by a

further developed and thicker “random structure” (cf. Shibata [ 150,151 ). Qualitatively, they

noted one or another of the following processes as being rate-determining:

i) electrosotption of 0 or OH “ions” (32,78.40,55,133];

ii) removal of protons and desorption of H,O molecules: and

iii) place-exchange between Pt and OH or 0 species [78,51,152].

Extents of 0 coverage (8,) were based on the ratio 0,=QJ2Q, for the charges Q, for oxide film

reduction and Q, for H coverage, 8,=1. 8,=1 they took as corresponding to an oxygen

monolayer of 1.3 x IO” atoms per real cm*, representing, they stated, MI oxide layers of Pt(II)

oxide, in a place-exchanged configuration. This is the state we refer to elsewhere (Sections II.4

and 11.5) as the quasi-2-d state of oxide film formation. It appears from this paper that Vetter and

Schultze did not distinguish the possibility of the first layer being OH chemisorbed on Pt with a

second stage of oxidation and de-protonation to 0, with place-exchange, though there are various

indications that it is an “OH” state that is fust e lectrodeposited, with subsequent Pt/OH + Pt/O +

H’ + e conversion.

Although these authors wrote the formal process of oxidation at Pt in terms of the reaction Pt +

H,O + Pt - 0 + 2H’ + 2e, they regarded the oxide film forming process as involving transferred

ions viz. H,O (aq.) + Oad2- 2H’ (aq.). Presumably the 2e must enter the Pt conduction band to

account for the coulombically measurable charge that is seen to be passed in cyclic voltammetry

or galvanostatic pulse experiments. However, it is possible that O*- ions (or OH- ions) are

involved since the place-exchange process, which produces ultimately a quasi-bulk type oxide

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film, is believed to take place (initially) between anodically generated PI? and the O*- (or OH-)

ions. This matter is difficult to settle since the charge-state of adsorbed species at the Pt interface

is unknown and 0” or OH- may be chemisorbed with partial charge transfer (cf. the

electrosorption valence [60,61]), as found by Angerstein-Kozlowska et al. [56,57] in their detailedstudy of the microscopic process involved in the initial stages of oxide film formation at Au.

(See Part III which follows). Ultimately, this question revolves around the extent to which thin

oxide films at Pt or Au are ionic or covalent [ 1561.

An additional complication that received little attention in earlier work was that the deposition

or chemisorption of OH or 0 species at Pt and also Au [45,56,57,153] is directly coupled with

desorption of anions of the electrolyte that are already pre-adsorbed before onset of oxide film

formation takes place. However, this factor was already indicated by the works of Frumkin[ 1551, Balashova [62], and more recently by Horanyi et al. [138] and Wiechowski [ 1541. It is an

important aspect of the mechanism of oxide film formation in its very initial stages (see Section

1.9) and complicates the simple reaction scheme considered by Vetter and Schultze.

From cathodic galvanostatic charging transients, they showed (Fig. 22) that the oxide formed (as

charge Q, or 0,) increased linearly with the potential of formation, in the range 1.0 to 2.0 V, RHE

but with a decreasing slope for decreasing formation times in the range 1000s down to 2~10.~ s.

By commencing their experiments at 1.0 V, Vetter and Schultze evidently missed the interesting,reversibly formed, submonolayer of OH species which can be resolved in cyclic voltammetry

Mow that ootential 1551.

1.0

b)I , I J

0.8 1.0 1.20

Potential/V, RHE

Fig. 22 a) Linear relations for increase of extent of oxide film formation (Q) at Pt withpotential in the range 1.0 to 2.0 V(RHE). (From Vetter and Schultze, ref. 58)b) Linear relation for QOHvs potential from linear potential sweep experiments,

showing inflection at monolayer coverage at cu. l.OV, RHE. (From ref. 55).

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378 B.E. Conway

Their galvanostatic reduction transient data enabled them to show that the previous growth of

the oxide film was linearly logarithmic with time of formation at a given potential, confirming

the earlier observations of Gilroy and Conway [142], Gilroy [143] and supporting the later

mechanism of Conway et al. [ 1571. The slopes of the growth relations, dQJd log t increased with

formation potential, as also noted previously [ 143,157].

They also deduced that the oxide developed continuously with (log) time, contrary to the

conclusions of Biegler and Woods [121] and of Parsons and Visscher [70]. However, this

contradiction was resolved in recent work of Conway et al. [158] (see Section ITS) who showed

that it was only the quasi-2-d state of Pt oxide formation that reaches a saturation limit, not the

overall extent of oxide film formation which can continue to multilayer extents of oxide film

development [ 150,151,159] in a logarithmic way in time.

The direct logarithmic law of oxide growth (- Q), they showed, corresponded (as it has to) to

an anodic growth current inversely proportional to time (/i.dt = j(Wt)dt). Some complications can

arise in the growth experiments conducted above ca. 1.7 V, RHE, due to unavoidable

co-evolution of 0, which can then be reduced in the cathodic transient. This effect will increase,

of course, at higher polarization potentials. From their data Vetter and Schultze were able to

construct isosteric current vs potential curves over the range 8, = 0.25 to 2.5, the latter region

being for higher potentials and longer growth times.

Two other interesting aspects of their and others’ results were as follows:

i) Provision of evidence that an anodically formed oxide film at Pt underwent a small but

significant ageing effect on open-circuit (also found by Arvia), leading to a greater stability to

reduction (less positive reduction potential, like that required for reduction from an higher

coverage). This ageing effect is attributable to continuing rearrangement of the oxide film after

its initial state of formation.

ii) An indication by Dietz and Giihr [ 1601 that the reduction of thick films is non-uniform, i.e. it

proceeds at the edges of oxide islands (cf. ref. 95), as was deduced by Schultze and Vetter [161]

in their related work on oxide film formation and reduction at Au [45]. The charging curves had

the same shape as complete curves on a reduced time or charge scale which Vetter and Schultze

deduced as corresponding to the requirements of the “island” reduction mechanism (cf. STM

images, ref. 244) first treated by Conway and Marincic 1951.

From their galvanostatic transient experiments, they were able to construct Tafel relations both

for oxide formation and oxide reduction over substantial ranges of coverage (0, in anodic

polarization, 8, in cathodic) as shown in Fig. 23.

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is (oxide rrduct:onl

1r+(oxide formation)

A.nnn , r

h\

0.8 10 1.2 1.4 1.6 18 20

potentiol cH/V -

Fig. 23 Anodic and cathodic Tafel relations for Pt surface oxide formation and reductiondeduced from respective transient response measurements (from Vetter andSchultze, ref. 59).

t 120.

1onodfcfb,)

5100 -

41280

=o 60 -

:40- 5 “eL 0 0

-2 -0C-0

0 1.0 2.0 0 IIIztr1.0

8-overage

Fig. 24 Tafel slopes b, and b. for oxide film formation and reduction, respectively, at Pt.(From Vetter and Schultze, ref. 59).

The anodic and cathodic Tafel (b) slopes (b, and b., respectively) were quite different (Fig. 24)

and b, increased linearly with coverage [b, = 36 (1 + 1.08) mV]; b. was cu. 60 mV, more

independent of 8 except at 8 < 0.8.

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380 B.E. Conway

The cathodic reduction currents are a function of 0, but depend on pre-history conditions for the

development of the film. Under some conditions, the log [cathodic current] increases with 0,

corresponding to a relation demonstrated by Bagotzki et al. [162].

Based on considerations of interfacial capacitance, the model of the oxidized metal interface,

shown in Fig. 25, was used as a basis for interpretation of their results on oxide growth kinetics.

Essentially, this description differs from the Stem-Grahame one only in respect of the

introduction of the growing oxide film between the metal and the inner #Helmholtz plane (Fig.

25) and a significant potential drop across it allowing electron tunnelling. When the film is very

thin, viz. a monolayer, the potential drop at the surface has the nature of a surface dipole

potential difference [ 1571, coupled with an external ionic double-layer. There will be an internal

potential drop at the metal/oxide interface, not included in the diagram of Vetter and Schultze.

Fig. 25 Multi-interface capacitance model for oxide formation at noble metals andcorresponding potential (t$) profile. (From Vetter and Schultze, ref. 58).

Apart from its very initial stage, Vetter and Schultze considered the oxide growth kinetics to be

controlled by the rate of place-exchange between 0 ions and Pt cations generated from the Pt

metal surface, determined by rates of displacement of the ions and vacancies. The current-density

for Pt ion migration and oxide formation at the oxide/solution interface was written [58,59] as

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i+ = k, exp 5: .(A+z+ + A-z-)d

& (4% - 4,)0.x 1 (31,

where (I$, - I&)&,,, the metal to film field, could be replaced by (t$r - $,&Id in which 2 is a

parameter relating the interfacial capacity C, to the inner-layer capacity Ci (see Fig. 25) by (cf.

refs. 133, 163 and 164):

C;’ = C;’ (1 + a@ (32)

with a = dD,/S,D,,. Here d is the reduced film thickness &,,/e, Si the thickness of the inner

Helmholtz layer, and Di and D,, are the dielectric constants of the respective regions of the

interphase. The h quantities define the fractional jump distances h+d and X.d for positive and

negative ion displacements, as for high-field growth [ 1171.

Closely following the papers of Vetter and Schultze [58,59] was the significant paper of

Angerstein-Kozlowska et al. [55] on a detailed study of the early and later stages of Pt surface

oxide formation (and reduction) by means of linear potential-sweep voltammetry and “potential

holding” experiments. Several features of their results are considered below.

Employing high-purity conditions, especially the use of pyro-distilled water [3 11, hese authors

were able to resolve the formation of the initial 2-d oxide film state (which they attributed to

electrosorption of OH species up to a monolayer at cu. 1.1 V, RHE) into three

distinguishable states at polycrystalline Pt. They attributed these resolved states to the

progressive formation of overlay sub-lattices of OH on Pt. (see Section I.9 and Fig. 8).

While 2-d lattice models of chemisorption had been used in surface science to represent states

of chemisorption on ordered surfaces, it appears that, in their paper, viz. [55], this was the first

application in electrochemical surface science. Later, this kind of representation became

common, especially for describing the phenomenon of underpotential deposition of H and base

metals at noble metals where multiple peak currents in cyclic voltammetry of sub-monolayer

deposition and desorption processes arise even at single-crystals, in fact with better resolution

than at poly-crystals, e.g. as in the careful UPD study of Pb ad-atom deposition on the

principal-index planes of Au by Engelsman et al. [90].

Angerstein-Kozlowska et al. [55] made a quantitative resolution of the monolayer current profile

for OH deposition at Pt in linear-sweep voltammetry in terms of sub-lattice coverages in the

resolved peaks (designated as OAl, OA2 and OA3) in Fig. 4, as recorded in Table 2.

Fig. 22b also shows the integrated oxide formation charge as a function of potential; it has a

significant inflection at qO$qH = 1 at a potential of 1.1 V which was attributed to completion ofthe OH monolayer (“Pt,OH”). Beyond the latter stage, further development of the oxide takes

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382 B.E.Conway

place through the surface reaction: Pt, /OH + Pt 10. + H’ e, as in Scheme 1 (Section I. 14).

Table 2

Region Charge to peak/ PeakpC cm-* potential/V 8,, (based on 1e/Pt atom)

ON0

A2

0

Biiad regionbeyond 1.2 V

26-35 0.89 0.12-0.16

81-88 0.94-0.95 0.37-0.40

172 1.04- 1.05 0.78

1.1 v 1.01.2 v 1.351.3 V 1.651.37v 1.9

The charges associated with formation of the successive states of OH occupying the (100)

surface lattice are as shown in Table 3. The occupation geometries were illustrated in Fig. 5 for

the (unreconstructed) (111) and (100) lattices.

Table 3

Structure

Pt,OHP&OHPtOHPto

Charge per structure

pC cm*

55

55

110220

Total charge

pC cm*

55

110220

440

Charge at peaks

pC cm*

27.5

82.5

165

Broad region

The basis of the sub-monolayer resolution is illustrated in Fig. 5 and Table 4 where the

observed pseudocapacitances for the three peaks (OA 1, OA2, OA3) are compared with the

theoretical values expected for the three sub-lattice states according to the following surface

lattice occupation processes:

Pt + H,O + Pt,OH + H+ + e (Q = 55 pC cm-*) (33)

Pt,OH + H,O -+ PhOH + H+ + e (Q = 55 pC cm-*) (34)

Pt,OH + H,O + Pt,OH + H+ + e (Q = 110 pC cm-‘) (35)

Resolution of such stages of oxidation of the surface appears to be a general phenomenon since

quite similar behavior is observed at Au where the very initial stages of oxidation can also be

followed (see Part III). A good computer simulation of the anodic i-V profiles can be made on

the basis of the above processes [172].

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Table 4 Peak Pseudocapacitances for OH Electrosorption Stages in Comparison with TheoreticalValues for the (111) Pt Surfaces

C&F cm.*

Peak Expt. (111) (100)

Individualpeaks

AE, =75 mV

Individualpeaks

AE, =

75 mV

O*, 660-750 783 920 510 60004, X00-X60 7x3 1040 510 6770 A? 792-830 783 1020

In addition to processes 33 to 35, further oxidation of the surface beyond monolayer “PtOH”

occurs with increasing anodic potential and with holding the potential at various values above

0.85 V. The further oxidation process must be (cf. refs. 5X and 59)

PtOH 3 PtO + H’ + e (q, = 220 pC cm-‘: total qa = 440 pC cm-*) (36)

or thickening of a rearranged PtOH layer. Process (36) occurs principally over the broad region

(Fig. 4).

As the monolayer of OH species is developed towards the limit of nominal full coverage

(“Pt,OH”), it is believed that the film becomes progressively place-exchanged (Pt/OH + OWPt)

commencing already around 8,, = 0.15 (at 298K) and depending on anion co-adsorption. This

accounts for the progressive increase of irreversibility between the processes of oxide film

formation and reduction. The detailed examination of the latter phenomenon was the second

principal contribution in the paper of Angerstein-Kozlowska et al. [55].

11.2 Reversibility and Irreversibility of Various Stages of Surface Oxidation

i) Effect of anodic reversal potential E, in the sweep. The different ranges of potential over

which the formation and reduction of surface oxides occur on the noble metals is a striking

feature of the electrochemical behavior of these metals. It is also reflected [ 165,166,167] in a

kinetic irreversibility of most electrocatalytic oxidation reactions that proceed on these metals due

to the different range of potentials over which the surface oxide exists in anodic and cathodic

directions of change of potential. The irreversibility in formation and reduction of surface oxide

can be characterized in two ways: (a) in terms of different regions of potential over which the

oxide is formed and reduced and (b) in terms of the electrochemical kinetics of the processes

involved.

An overall i vs V profile taken to say 1.4, 1.2 V and reversed, shows behavior typical of

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384 B.E. Conway

hysteresis: e.g. (as in Figs. 17,26), both the potential range of the reduction and the shape of the

reduction peak are different from those for oxidation. An irreversible process in the anodic

sweep is indicated by the observed continuation of anodic current in the initial part of the

cathociic sweep. Potential sweeps taken to progressively less positive potentials show that the

electrosorption of oxide material becomes less irreversible (Fig. 26). Thus, anodic-going sweeps

taken up to a potential of only E, = 0.85 V in the O,, region (curves l-3 in Fig. 26) give, upon

reversal, an immediate fall of the current and a corresponding cathodic peak at almost the same

potential in the reverse direction of sweep. This is the situation that at low temperature, enables

an envelope of reversible current profiles to be distinguished (Fig. 13).

At 298K, irreversibility already begins to set in between the formation and reduction of the

species formed up to a potential of 0.95 V, i.e. well below the monolayer limit of 1 OH per Pt at

I. 1 V. For the remaining curves obtained for polarizations to successively higher E, values,

irreversibility increases, manifesting itself as a progressively greater difference between the shapes

of the anodic- and cathodic-going curves (Fig. 26).

I I0.5 1.0

potential (RHE)/V

Fig. 26 Progressive increase of irreversibility between oxide formation and reductionprofiles at Pt with oxidation taken to increasing positive limits in the anodic sweep.

An important question arises as to whether the separation between the anodic and cathode i vs

V profiles of Figs. 4 or 26 arises for kinetic reasons or because of true hysteresis (Scheme 1)

[78,140,141] involving an irreversible transformation of the initially electrosorbed OH species into

another species, e.g. by place-exchange, reducible only at less positive potentials than those

required for its deposition. The kinetic aspects of this question will be considered next.

ii) Kinetic significance of the shape of the i-V profiles. An important aspect of the behavior

at E, > cu. 1.0 V is that, upon reversal of the sweep the current does not immediately become a

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(mirror image) cathodic one as it does for a reversible process in the H or dl regions, but anodic

current persists in the cathodic sweep, as mentioned earlier, until the potential has become

significantly less positive and eventually, low enough for reduction to commence (see family of

curves in Fig. 26). Anodic oxidation of the surface thus continues during the initial stages of

cathodic sweeps for E, > 1.0 V. This is also supported by ellipsometric results where A continues

to increase briefly after dV/dt has changed sign in a sweep. This behavior is characteristic of an

irreversible oxidation process that requires overpotential for passage of significant current (cf. the

results of Vetter and Schultze, ref. 58). For an ordinary continuous charge-transfer controlled

Faradaic process, driven in anodic-or cathodic-going directions, the i-V profiles would be parallel,

with the two branches separated at all potentials by twice the double-layer charging current. For

an irreversible surface oxidation process, however, coverage is being progressively increased to a

limiting value LO hat the cathodic-going profile after reversal of the sweep will lie always below,

i.e. at lower anodic current values, than the anodic-going profile.

Another aspect of the irreversibility in reduction is that if a cathodic potential sweep is arrested

for a short time at a potential positive to the oxide reduction peak potential, the current at that

potential will slowly decrease; upon re-establishment of the cathodic sweep it is found that a

significant or appreciable decrease of charge for the remaining reduction is required. Given

enough time, almost all the oxide can be reduced at a potential significantly positive to the peak

potential. In fact, this extent of reduction on potential holding is logarithmic in time, as is the

potentiostatic growth. An example is shown in Fig. 27 which also indicates that two stages in the

remaining reduction can arise at Pt, as also found for Au.

Fig. 27 Two-stage reduction of Pt oxide film in a cathodic sweep taken from a previousanodic polarisation to 1.5 V.

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366 BE. Conway

Similar behavior was found by Gilroy and Conway [142] and by Tikhomirova et al. [168] and

by Will and Knorr [30], as well as in the work of Shibata [200].

iii) Kinetics of the anodic process. Kinetic aspects of a surface process can be followed by

means of response to a galvanostatic [58,59], a potentiostatic step modulation [22] or in linear

potential-sweep experiments [79] over a wide range of sweep-rates. When such experiments are

conducted in the anodic direction at Pt the peak potentials for the peaks OA,, O,, and O,, depend

very little (lo-20 mV) or not at all, on sweep rate in the range 0.005250 V s-l, indicating [79]

that the oxidation stages corresponding to these peaks behave in a kinetically facile way.

According to the treatment of Angerstein-Kozlowska and Conway [79], such stages have high

rate constants corresponding to large s, values, the latter representing (like i, for a regular

continuous Faradaic process) the Limit in (anodic) sweep-rate, s, below which the process remains

reversible, i.e. the peak potentials, Er, are independent of log s or s. Beyond s=s,, E, increases

logarithmically with s, as for a Tafel relation in log i. The peak potentials for processes OAl and

OA2 are virtually independent of log s over 2% decades of s and the E,‘s for the OA3 process

are also almost independent of log s.

This is an important and significant result since it means that the 2-d electrosorption of OH and

0 species in the short-time regime is a fast process. This behavior may seem at variance with the

observations [59,142,143,157] that growth of the oxide film at Pt, and also at other metals, is

usually a slow process, logarithmic in time. However. it has an essentially simple explanation,

and one fundamental to the understanding of oxide film formation at noble metals, as follows.

An essentially 2-d process involving electrosorption according to a rate equation of the form

can be rapid until 8,, for the sub-lattice involved approaches its limit; this can apply to the other

lattice occupancy stages OA2 and OA3 (Fig. 5). The observed behavior up to ca.0.95 V

corresponds to the above backward and forward partial processes taking place at comparable rates

so that E,‘s are more or less independent of the linear rate of change of E with time. The

long-time effects must therefore originate for a different reason [39,78]: this is that another

post-electrochemical process sets in allowing more “coverage” by OH-O species to be deposited

at a given potential than corresponds to the equilibrium quantity determined by reversibility in the

2-d processes, .e. by the Frumkin (or Langmiur, g=O) types of isotherm:

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for each of the sub-monolayer lattice-occupying processes referred to earlier. This slow

post-electiosorption process is believed to be the place-exchange reaction which allows slow

electrodeposition of OH and 0 species to continue beyond the equilibrium coverage condition, for

any potential, determined by the electrosorption isotherm above. In fact, this latter process

continues to post-monolayer coverages where more than one nominal monolayer of OH or 0

species is deposited, as envisaged in the representation shown by Vetter and Schultze.

This post-electrosorption type of process seems to be general as it is observed at all the noble

metals and leads to the direct logarithmic law of oxide growth at a given potential, beyond the

level of rapid formation of the initial, 2-d state from which the place-exchanged state and further

growth originate. It was illustrated in Scheme 1 earlier.

The initially electrosorbed OH (Fig. 16) appears to be metastable with respect to a rearranged

state so that, upon conversion to the latter, a more cathodic reduction potential is always

exhibited. These are the essential conditions [140,141] for the observed hysteresis. It seems (see

Fig. 26 and 13) that only quite a small coverage by OH species is required for some of the

species to become rearranged and the larger the coverage (higher E,) the greater is the extent of

rearrangement, giving a larger peak at the more cathodic potentials. In a rearranged Pt/OH

surface layer there will, however, always (see Fig. 16) be some OH species which are similar in

state to those initially electrosorbed and it seems reasonable to suppose that this fraction could

account for the fairly constant quantity of reversibly reducible OH species that can be

experimentally observed.

The driving force for rearrangement amongst OH or 0 species and Pt atoms will not only be

stabilization by change of the sign of the “Pt-OH” or “Pt-0” surface dipoles in the initial

electrosorbed state, due to their interaction with the interfacial field, but also relief of lateral

surface dipole-dipole interaction. Initially, some lateral interaction with chemisorbed anions,

which will have the same sign as that of the surface dipole moment, provides an additional

driving force for place-exchange.

iv) Time effects in the irreversibility The slow component in Pt surface oxidation, referred to

above, can be investigated by examining the effects of holding the potential for various times z,

at several E, values at the end of the anodic-going sweep and then applying the cathodic sweep.

Figure 28 shows the results of such experiments. In reduction curves obtained from E, = 0.89 V

(Fig. 28a), there is initially an almost reversible peak but the peak potential becomes

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388 B.E. Conway

progressively less positive with increasing 2, up to 30 min. In Fig. 28b, with E, = 0.94 V, there

is again a fairly reversible peak but E,,,, progressively moves more cathodic with time and,

furthermore, Q, increases appreciably with z,. For E, = 1.0 V, Fig. 28c indicates already

appreciable irreversibility between the anodic formation and cathodic reduction processes; again

E,,, becomes more cathodic and Q, increases with Z, as measured by the charge under the

reduction curves. It is evident, for example in Figs. 28b,c, that although only the initial stages of

surface oxidation of the platinum are still involved, there is nevertheless an appreciable time

effect with regard to continuing growth and/or rearrangement of the surface species forming the

oxide film. Already, at submonolayer coverages, the increase of Q0 is logarithmic in ‘L’,

Fig. 28 Series of anodickathodic voltammograms at Pt showing progression ofirreversibility in reduction with potential holding at several potentials, E,. (Holdingat E, also increases quantity of oxide formed and therefore available for reduction).a) E, = 0.89V; b) Ea = 0.94.

These time effects in the anodic process, which affect the kinetics of the cathodic process, also

arise when the anodic region is traversed at different rates, i.e. during potential-sweeps at various

s values, especially when s is small, say ~25 mV s-r.

11.3 Kinetic Interpretation of Shapes of i-V Profiles for Formation and Reduction of

Oxide Films on Pt

It was explained earlier that, except for the first 5 to 15% of coverage of Pt anodes by OH or 0

species, the reduction of the film is always irreversible with respect to its formation no matter

how slow is the rate of sweep in CV experiments; in fact, in slow sweeps, the irreversibility is

greater due to more time being provided for place-exchange to take place, i.e. more hysteresis

[ 140,141] is observed.

Two attempts have been made to provide a reaction-kinetic type of explanation of the

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irreversibility: one by Appleby [ 1711 and the other by Tilak et al. [ 1721. For a general single-

step electrosorption process, Srinivasan and Gileadi [76] showed that with increasing sweep-rate,

beyond s=s, (cf. ref. 77), anodic and cathodic peaks progressively separate from one another on

the potential scale but have the same geometrical forms (an asymmetric single peak for SX,; a

symmetrical peak for sas,). The behavior of Pt oxide formation and reduction is quite different.

It is mechanistically of major significance that in the anodic direction of film formation, peak

potentials (in so far as they can be resolved) are almost independent of s up to cu. 100 V S’, yet

on reduction, the E, values are linear in log s, i.e. Tafel behavior applies, characteristic of kinetic

irreversibility.

Appleby [ 171] considered the 2-step reaction

Pt + H,O -+ Pt/OH + H’ + e (39)

Pt/OH + PtO + H’ + e (40)

and wrote the respective net-rate equations as

q,(de,/dt) = q, s (d&/dV) = k,F [(1-fI,-0,) {exp (l-p) VF/RT . exp -(l-p) g,(B,+B,-%)/RT) - 8,

{exp - PVF/RT, exp Pg,(B,+B,-%)/RT)] - q,s (d8JdV) (41)

and

q,(d&/dt) = q,s (d@JdV) = k,F [e,(exp (1-p’) F (V-V,)/RT . exp -(I-p’)g, (8,+ 8, - %)/RT) - 8,

( exp - PIF (V-V,)/RT . exp p’gz (8, + e,-%)I (42)

where qr is the charge for monolayer deposition of the oxide species, 8, and 8, are the coverages

by OH and 0, g, and g, are the Frumkin parameters for lateral interactions in the interphase

(probab!y a common g value is required as all interactions are mutual), s is the sweep-rate in V s-

I, V, is the (negative) potential difference between the standard potentials for reactions 39 and 40,

8=% is the standard state coverage, and k, and k, are the rate constants. The total current flowing

is q,s (de,/dV + 2 d&/dV).

Numerical kinetic simulations were made on the basis that in the cathodic direction, process (-

39) was rate-determining while in the anodic direction, at least at low 8, the step (39) is

sufficiently rapid to be regarded as at quasi-equilibrium (cf. ref. 79).

It is to be remarked that in the reaction scheme (39), (40) considered by Appleby, no post-

electrochemical step of place-exchange was included, i.e. no account is formally given of the fact

that the species, after its initial deposition, is converted into another state rom which the

rrduction takes place. However, it was stated that the assumed exponential fall of the rate

constant of the first step with coverage (through the g, factor) was associated with rearrangement

of the film.

An important proviso was noted: that in order that two separate peaks are not resolved

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390 BE. Conway

(experimentally, under some conditions they can be; see Fig. 28) in the cathodic sweep, the

standard potential for the second oxidation step (40) should be less positive than for step (39).

For that reason, and because of the required high k, value, reaction (40) should be in quasi-

equilibrium in the anodic direction. This is consistent with the experimental results of

Angerstein-Kozlowska et al. [55] who demonstrated an high s, value [79] for anodic sweeps.

Some of the general features of the cyclic voltammetry behavior of Pt oxidation/reduction were

reproduced by Appleby’s calculations, based on empirical assignments of the values of the

parameters of the treatment. In Fig. 29, below, we reproduce three of his simulation diagrams,

/ v(vo’ts)

Fig. 29 Simulation i vs V profiles for formation and reduction of an oxide film,reproducing qualitatively the irreversibility or hysteresis observed. (From Appleby,ref. 171, q.v. for parameters used).

A more detailed series of calculations were given by Tilak and Conway [ 1721, aking account of

the successive stages of sub-lattice occupancy [55] by OH species. They also considered the two-stage mechanism, represented above, but included the back-reactions and their rate constants in

the reaction-kinetic scheme. They also included a two-stage mechanism involving bridged 0

species

Pt/OH + Pt +Pt\

O+H++ePt’

(43)

that can lead to a self-inhibitory behavior. Their calculations gave a quite good representation of

the observed anodic current profile for Pt oxidation up to the Pt/O stage, including resolution of

the 3 sub-lattice OH peaks and the subsequent broader 0 electrosorption region.

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For the oxide reduction process, Tilak et al. [ 1721 deduced sets of conditions that would

account for observation of a single peak on reduction exhibiting a Tafel slope dEr,ld log s of cu.

120 mV. Based on reduction of Pt/O to 3 sub-lattice states [55], P t,/OH, PtjOH and Pt/OH, they

demonstrated how an overall single peak could be generated in reduction, exhibiting either a 120

or a 40 mV (as observed) slope dEdd log s. These authors also specifically considered the more

realistic situation that, in reduction, the cathodic processes take place from place-exchanged

states, (Scheme 1).

II.4 The Quasi-2-d State of Pt Oxide Film Formation as Revealed in Reduction

i) Phenomenology and relation to other states. An important complication in the formation

and reduction of Pt oxide films [55,78,172] was revealed in experiments by Shibata [ 150,195] and

by James [ 1791 which showed that upon extended anodization at relatively high potentials (2.0 -

2.2 V, RHE), one or more further states of the oxide film, in thicker layers (cf. Burke [ 180.26]),

were generated, characterized by reduction in peaks at substantially less positive potentials (0.2 to

0.45 V, RHE) than those for the OCl peak normally observed. Conway and Jerkiewicz [ 1591

found that up to three further states (or stages in the kinetics of reduction) could be distinguished,

depending on the conditions of anodic film formation. Thick-film states of oxide formation can

also be generated at Au, as found by Barnartt [63], and at Rh, by Burke [26].

Several directions of work have been devoted to characterization of the behavior of this quasi-

2-d film, designated OCl, in relation to other states that can be revealed. It is convenient to

identify these approaches as follows:

a) demonstration that the OCl state at Pt reaches a limit of extent of formation with increasing

time of anodic polarization [refs. 121,122 and Section 11.51.

b) demonstration that such a state can also he characterized optically by means of ellipsometry

in a way that also corresponds to a limit to its extent of formation [refs. 70, 198 and Section 11.51.

And

c) demonstration that, under conditions for which the thicker layer states can be generated

[ 159,195], the OCl state remained at its limit of formation (corresponding to a state nominally

represented by “PtOPtO” [ 119,122] and moreover could be reduced independently of the other,

thicker film states under linear potential-sweep conditions with appropriately chosen potential

limits (see Section 11.4.ii below).

In the following material, we review the results and conclusions concerning matters (a), (b) and

(c), distinguished above, that are of great importance for the full understanding of the

phenomenology and mechanism(s) of oxide film formation at Pt.

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392 B.E. Conway

ii) Independence of formation and reduction of 2-d oxide on Pt from presence of thicker,

phase-oxide layers. During the course of recent studies of formation and reduction of thick-film

phase-oxide layers, two independent but closely similar pieces of work have appeared that showed

that the quasi-2-d film can be independently formed and reduced in the presence of thick-film

oxide material at Pt electrodes: one by Farebrother et al. [197] and the other by Conway

et al. [ 1481”.

The matter originates from the conclusions of Shibata [ 150,174,195] regarding his results on

thick-film oxide (his “p” state) reduction. He deduced that reduction of the quasi-2-d oxide film

(his a state, our OCl state) in a negative-going potential sweep led to deposition of P t atoms on

the outsidr of the thicker, phase-oxide film. As indicated in a paper by James [179] (see

comments in ref. 179 and in James’s “note added in proof” in ref. 179), Shibata’s view of the

course of reduction of the quasi-2-d film and the thick film implied that the thick (p) oxide would

become sandwiched between the bulk Pt metal and a film of newly deposited Pt atoms on the

outside of the remaining p oxide, resulting from the reduction of the a state of Shibata. Burke

and Roche [180], in their paper on development of so-called thick hydrous oxide on Pt by a

cycling procedure, and Burke in his review [180], recognized a similar situation, opposing the

conclusion of Shibata [ 1741.

In neither of these papers, however, was any experimental evidence provided which could

distinguish these viewpoints. Here we summarise the recent experimental evidence [197,148]

which provides a rather clear indication in favor of the quasi-2-d state lying beneath that film and

being reducible independently of the then overlying thick oxide film. Settlement of this question

is important (a) for theories of growth of Pt oxide films and (b) for identifying the state of the

external surface of thick, anodically grown oxide films as electrocatalyst interfaces on which the

Cl, (cf. refs. 17,27,52) and 0, evolution reactions proceed [ 184,199].

First it is necessary to outline the behavior that is observed in the sequential reduction of

anodically formed states of P t surface oxide.

In Fig. 30, curve 1 illustrates the normal cyclic-voltammetric behavior for sweeps taken to 1.4

V, RHE from 0.05 V, showing the OCl state (Shibata’s cx oxide), well characterized in previous

works. Curve 2 of Fig. 30 shows the linear-sweep voltammogram for Pt surface oxide reduction

taken after formation of the quasi-3-d state (cf. ref. 174) at an elevated positive potential of 2.0

**During a presentation at a scientific meeting, these authors found that they had almostsimultaneously made the same kinds of experiments and reached the same conclusions, so theydecided to publish their findings independently, without further discussion at that time.

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V, RHE, with an arrest period being provided in the negative-going potential sweep at the zero-

current potential, E,=,, (cu. 1.70 V) to remove any anodically generated 0, by N2 bubbling before

the sweep is taken into the potential regions of oxide film reduction.

Fig. 30 Diagram of experimental cyclic voltammetry curves for successive formation andreduction of oxide films at Pt formed to the OCl state (in reduction) and up tohigher potentials with reversal after covering the OCl region (from ref. 148).

Under these conditions, it can be seen that the quasi-3-d oxide film is reduced in a single peak

(designated OC3*), well separated from OCl which has now increased in charge. This OC3 state

is reduced at potentials near or within the H UPD potential range. Figure 31 also shows further

development of states of the oxide film as manifested in the reduction curve for a film grown at

2.30 V. Now additionally (cf. ref. 174) a state OC2 is developed on the current profile of which

OCl reduction currents are discernible just as a shoulder. The state OC3 (Fig. 30) disappears but

is replaced by a narrower peak (designated OC4), shifted to less positive potentials in the H UPDregion. (A broad shoulder, BS, develops increasingly over the potential range 1.1 to 0.6 V as

potential and time of anodic polarization increase, Fig. 31).

The result of most significance is that concerned with the question (cf. refs. 179,148) whether

the reduction of state OCl gives rise to Pt atoms on the outside or on the inside (i.e. on the bulk

Pt substrate) of the thick film. This result is obtained by experiments on reoxidation of Pt in a

potential sweep after incomplete reduction of an oxidized Pt surface exhibiting phase-oxide states

*The designation of the reduction peaks, OC, (“oxide cathodic”) is in the order of theirappearance in the negative-going sweep.

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B.E. Conway

Fig. 31 Reduction behavior of a Pt oxide film grown at 2.3V showing development of OC2and OC4 states (from ref. 148).

The following sequence of anodic film formation and reduction operations provides an answer

to the above question. First, at an electrode on which had been formed the stable OC3 state of

the oxide film, (as revealed subsequently as curve 3 in Fig. 30), together with OCl (Fig. 30), the

negative-going sweep is reversed at a potential where reduction of the OCI state has just been

completed (cu. 0.33 V, RHE) and a following positive-going sweep is applied. Fig. 30 shows

that this sweep traces out a repetition of the anodic-current profile (shown as curve 3 in Fig. 30),

over the potential range 0.85 to cu. 1.4 V, that is characteristic (cf. ref. 55) of formation of the

quasi-2-d array surface oxide observed at an initially unoxidized Pt surface, yet in this case the

stable oxide state OC3 is evidently still present on the s&ace as is indicated by the reduction

peak observed on a continuation of the sweep after reversal to the negative-going direction again,

as shown in Fig. 30, as curve 3. Finally, when curve 3 continues again in the positive direction,

it traces out the oxidation current profile first for oxidation of UPD H, followed by redeposition

of the quasi-2-d OH and 0 states between 0.8 and 1.4 V, RHE. The sequences of potential

ranges in the sweeps are indicated in the annotations on Fig. 30 and in its caption.

The behavior described above suggests that the 2-d surface oxide formation and reduction

processes are independent from those that lead to formation of the more stable, thicker films,designated OC2, OC3 or OC4 here. As was mentioned, Shibata [174] suggested that reduction of

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the OC I state (referred to as “a” in his paper) leads to a Pt atom film residing on top of the more

stable oxide film OC3 (denoted “p” in his paper). However, from Fig. 30, it is clear that the

details of the anodic current profile, following reduction of the OCl state, correspond closely to

those observed [30,55] on a bulk-metal Pt surface; also such features are now known to be

characteristic of the crystal surface orientation in the case of single-crystal Pt electrodes. It seems

unlikely that such details would be reproduced on reoxidation of a “monolayer” of Pt atoms

residing on the outside of the remaining stable hydrous oxide film.

These results therefore indicate that the thick-film layer of oxide states, designated as OC2, OC3

and OC4 here, are microscopically porous, as concluded by Burke [ 1801, and enable the quasi-2-d

oxide film to be regenerated, after reduction of the OCl state, underneath these quasi-3-d.

thicker-tilm states which evidently still remain on the Pt surface, and independently of their

presence. Correspondingly, since the OC1 reduction peak is always separately observed, even

after strong anodization which generates the OC2, OC3 and OC4 states as well, it seems that

some compact quasi-2-d state of the oxide must always remain in contact with the underlying Pt

metal surface and must have a quite different structure and stability from those of the quasi-3-d

states that can be developed at high potentials and long times of film growth. A similar

conclusion was reached by Allen et al. [ 1941 rom XPS measurements that “PtO,” resides on top

of “PtO”.

The behavior observed further suggests that, independent of the development of the 3-d phase

type oxide states, the underlying metallic surface of the bulk Pt always programs the formation of

a kind of 2-d array state immediately in contact with the Pt. Thick oxide film growth must then

presumably take place by field-assisted transfer of Pt cations from, or through this film, into the

3-d type layer where the Mott-Cabrera (201 high-field growth mechanism can apply when the

resistive “p-state” (cf. ref. 180) is established. The field arises, of course, on account of the

imposed anodic polarization with a component associated with the surface Pt-0 dipoles [ 1.571.

We end this section by repeating that very similar conclusions were reached simultaneously by

Birss et al. [ 1971, published almost at the same time as the note by Conway et al. [ 1481. An

important result in the work of Birss et al. [ 1971 s that it was found that the thick film states

which normally seem to be reduced in the potential range 0.4 to 0.1 V, RHE, as cathodic current

peaks (OC2-OC4), could be reduced at substantially more positive potentials provided sz@cient

time was allowed, i.e. the peak potentials were determined to a substantial extent by the kinetics

of reduction which hence, in part, determines the hysteresis.

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396 B.E. Conway

II.5 Apparent Limit of Anodic Oxide Film Formation at Pt

i) Nature of the problem. In the course of researches on Pt oxide film formation and

reduction which were active in the late ‘60’s, Biegler and Woods [I211 and Biegler et al. [173]

showed interesting results that seemed to indicate attainment of a limit to surface oxide film

formation at Pt. Similar observations were reported by Visscher and Devanathan [152], Parsons

and Visscher [70], based on ellipsometry (see Section 1.12), and by Visscher and Blijlevens [ 1751.

Almost simultaneously, the same conclusion was reached by Russian researchers [ 176.1771. The

results reported in refs. 70,175 were based on ellipsometry measurements and were consistent

with the “non-optical” data reported in refs. 121,152. The “history” of the work indicating this

attainment of limiting oxygen (oxide) coverage, and discrepancies between various research

laboratories, is discussed in detail by Woods in a review chapter [ 1781.

The results of Shibata [ 1501, James [ 1791, and of Gilroy and Conway [ 1421 showed, contrary to

the indications of attainment of a limiting extent of oxide formation, that cnntinuous anodic

growth of Pt oxide to relatively thick films could occur, in fact in a logarithmic way with time at

constant potential. Inconsistencies amongst results of various researchers and problems of

reproducibility found in some works [ 180,18 ] were attributed to differences of electrode

preparation and pretreatment conditions which led to different oxidizability of Pt [ 178,180]. For

example, Biegler and Woods [ 1211 hemselves recognized that more extensive oxide film

formation could “sometimes occur” under severe conditions of anodic polarization (e.g. 1000 s at

2.98 V) but such behavior, they stated, was only occasionally and irreproducibly observed. On

the other hand, in work of the present author, continuously extended oxide film formation was

never a problem to observe reproducibly.

In the following paragraphs we examine the origin of this apparently contradictory situation. Its

clarification is quite important for a) theories and models of oxide film formation and growth at

noble metals and b) characterization of the external surfaces of oxide films on noble metals where

the oxide/solution interface is the electrocatalytic surface for various important electrode

reactions.

In considering the structure of the Pt oxide layer that had reached supposed limiting coverage

(but cf. ref. 158), Parsons and Visscher [70] suggested a “PtO” structure consisting of two

alternating planes: Pt-0-Pt-O/solution. A similar conclusion was reached by Jerkiewicz and

Conway [ 1581 who examined the significance of the supposed limit of oxygen coverage in

relation to the demonstrable development of much thicker films that are formed at higher

potentials and for long anodic polarization times as first noted by James [ 1791 and by Shibata

[ 150.1741. The “Pt-0-Pt-0” film corresponds to Pt in oxidation state +I1 as indicated by ESCA

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measurements [186]. The thicker films formed at higher potentials (see Section Il.4) are in

oxidation state +IV [ 1861.

Visscher, with Blijlevens [ 1751, extended her work with Parsons on the apparent limiting

coverage by oxygen species at Pt, found by Biegler and Woods [121]. There was some confusion

about the significance of this coverage limit as Biegler and Woods had referred to it as a

monolayer but having a charge ratio, qd2qn, referred to the charge for formation or desorption of

H, qri, of 2.68 which would not correspond to an actual monolayer of 0. However, in a later

paper, Biegler, Rand and Woods (1731 showed that the method of calculation of q, for a

monolayer of UPD H was incorrect, giving a revised value for qd2qr, of 2.05 which is much

nearer to the value for a true monolayer of “PtO” (one 0 atom per Pt atom in the surface, not a

stoichiometric compound). However, Tyurin and Volodin [ 1881 reported a limiting coverage

factor of 2.56 from cathodic sweeps at 40 V s-l which may have been influenced by some 0,

reduction unless the solutions had been purged with N, after anodic film formation (cf. ref. 159).

In the work by Visscher and Blijlevens [ 1751, coulometric measurements of oxygen coverage at

Pt were made at high anodic potentials, up to 2.6 V, in aqueous H,SO, over the concentration

range 1 to 4M. The apparent limiting coverages were found to arise above cu. 2.1 V but

depended appreciably on the H,SO, concentration diminishing from values (corresponding to

qJ2q,) of 2.20 down to 1.15. This result was attributed to chemisorption of some species derived

from H,SO, (HSO,- or a charge-transferred radical). Some Pt dissolution can also occur [ 1871.

Parsons and Visscher [70], in considering their oxygen coverage results for high potentials,

suggested that the oxygen species deposited below 1.5 V becomes “buried” at the higher

potentials (cf. Schuldiner’s “dermasorbed” 0 and the discussion in Section II.4 on the relation

between the 2-d state of electrosorbed oxygen and the thick film states). Then the results of

Visscher and Blijlevens would be explained by this deeper-seated oxygen being unaffected by

HSO,- but by HSO,- replacing the “outer” oxygen layer, thus giving rise to the diminishedlimiting oxygen coverage at high H,SO, concentrations.

In a recent paper by Jerkiewicz, Tremiliosi-Filho and Conway [ 1581, t was found that the

charge associated with the OCl state, in reduction, can attain values even over 2000 pC cm-2 (P

ca. 4.5 O’s per Pt). However, extended periods of oxidation lead to a complication in so far as

formation of relative thick films, upon reduction, generates increased real-to-apparent surface area

ratios, R, (cf. ref 121). The important consequence of this is that the apparent OCl reduction

charge values, Qorl, must be corrected for these increases of R; when this is done (Fig. 32), theQ,,,-, remains essentially constant and close to a value of ea. 880 pC per real cm2, independent of

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398 B.E. Conway

E,, in the range 1.80 to 2.30 V, RHE, and independent of the increasing quantities of oxide in the

other, electrochemically more stable states designed OC2, OC3, OC4, which are also formed

given sufficient time of polarization above about 1.8 V, RHE.

Fig. 32 Development of the OCl state before and after correction for real area change ofthe Pt substrate due to thick-film reduction (from ref. 158).

The states of hydration of such possible thick-film Pt oxide species were extensively discussed

by Burke et al. [ 1801 and other researchers [ 1951961.

It is believed that the OCl state is a quasi-2-dimensional compact film [55,190] within whichthe Pt is in the +I1 oxidation state while in the quasi-3-dimensional oxide (bulk-type oxide),

developed on strong anodic polarization, it is in the +IV oxidation state [ 186,192- 1941. These

results suggest the above Qoo charge of 880 pC cmm2 orresponds to a limit of 2 equivalent

monolayers of “PtO”. This is consistent with the conclusion of Parsons and Visscher [70] from

ellipsometry measurements that a limit of oxide film formation (corresponding to our OCl state)

is attained corresponding, in their representation, to “Pt-0-Pt-0”.

The observations of Jerkiewicz et al. [ 1481 clarify the somewhat puzzling aspect of earlierresults [70,121,175-1781 in the literature: thus it is evidently only the OCI state that reaches a

limit corresponding to a reduction charge of cu. 880 pC pr real cm2 while oxide film formation in

the other states (Figs. 30 and 31) continues without reaching any limit. Like other researchers

[121,178,148], Balej and Spalek [181] concluded that an overall limit (2.04 0 per Pt) was

attained but their conclusions referred to not very well defined surfaces or electrode pretreatment

conditions. Biegler and Woods [ 1211 measured an attained limit corresponding to 2.7 0 atoms

per Pt which they referred to as a “monolayer” of 0 on Pt. This obviously does not correspond

to a monolayer nor to the limit “PtOPtO” referred to by Parsons and Visscher [70]. However, if

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Biegler and Woods [121] had corrected for possible real-area changes, a real limit of a two 0

atoms per Pt would be found corresponding to the corrected 880 uC cm-’ reduction charge limit.

The development of this quasi-2-d oxide film up to the above definite limit of formation must

be considered in the light of the overall continuous growth of the film that is observed to be

logarithmic in time. In the following section, this matter is treated in more detail.

11.6 Kinetics of the Oxide Film Growth Processes

Continuous film extension in the states OC2, OC3, OC4 takes place under appropriate & and t,

values in parallel with formation of the OCl state up to its limit. The kinetics of this film

extension are of complementary interest and have been dealt with in some ways in various

publications [ 143,157]. Fig. 33 shows the oxide growth plots in log t,, (cf. refs. 142,143,157] for

polarization potentials between 180 and 2.30 V, RHE (for detailed discussion see refs.

[ 142,143,157,139]).

Fig. 33 Oxide film growth plots in log t,, for Pt in aq. H,SO, (298K) at holding potentials,

E,.

The main points to be stressed here are: (a) a direct logarithmic growth law applies

[ 142,143,157] and extends to E, values between 1.80 and 2.30 V; since the OCl state reaches a

limit (although, initially, it also increases with log th), it is, however, implicit that the overall

growth cannot be exactly logarithmic in t,,; (b) at sufficiently long t,, the oxide growth rate

drastically increases (some 400 times) due to the onset of a new mechanism, probably the “high-

field mechanism” associated with appearance of the much more resistive film (the “B” state

demonstrated by Shibata [ 1951) and (c) this drastically increased growth rate arises when an oxide

film has been formed to an extent of cu. 1800 uC cm-*, which corresponds approximately to the

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400 B.E. Conway

charge required for onset of high resistivity of the film [ 19.51.

Regarding point (c), the following explanation may be suggested: initially, the oxide is formed

only in the OCl state corresponding formally to 2 equivalent monolayers of “PtO” (880 pC cm-‘).

Subsequently, the formation of one nominal monolayer of “PtO,” (another 880 pC cmW2)n top of

two nominal monolayers of “P tO” is indicated (Fig. 34). These growth processes are slow

(logarithmic in time) but once they have been completed, the growth rate dramatically increases.

The total charge equivalent to 2 monolayers of “PtO” and 1 of “PtO,” corresponds to 1760 uC

cm-’ which is close to the observed value of 1800 pC cm-2 referred to above, beyond which rapid

film growth sets in. Here, note the reservation about stoichiometric representations of thin oxide

films made in Section 1.5. It seems that is it this state of the oxide at Pt that is required for good

efficiencies of anodic coupling reactions, such as the Kolbe and S,O,*- syntheses, to arise.

Under conditions of oxide film formation above room temperature at 323, 333 and 348 K in a

single sweep, it is significant that the OCI charge limit is still found to be maintained at 880 pC

cmm2 et the other distinguishable states continue to be generated but at higher rates at elevated

temperatures. These results confii the special significance of the 2-d OCl film as an

independently developed state of the oxide film at Pt having its own particular properties.

II.7 Model of the Thin and Thick Oxide Film

The experimental results reviewed above and described in the literature [ 180,1951981 now

allow a rather definitive model of the thin and thick oxide films on Pt to be proposed, as

illustrated in Fig. 34. The bulk-type oxide (“PtO,“) resides on top of the quasi-Zdimensional

oxide (“PtO”), clear evidence for which was given in refs. 190,194 and 197, as discussed in an

earlier Section, II.4.ii.

The quasi-2-dimensional oxide (OCl state) can be sepurutely reduced, most likely because the

overlying film is hydrous (cf. refs. 180 and 197) and porous, allowing transport of protons

through it [ 1801. This behavior, and the attainment of a limit [ 121,122] to the extent to which the

2-d film is formed, is consistent with it being in some qualitatively and observably different state

from that of the 3-d film.

The situation described here implies that the distinguishable thick-film states (OC2-OC4) are

formed either by injection of Pt ions thro~~gh he OCI film, leading to a film-forming reaction

producing the OC2-OC4 states, outside the quasi-2-d film or the OCl “PtO” material is

continuously formed at the Pt metal surface but converted, at its external interface, to the “PtO,”

material which grows without apparent limit. There are difficulties with the latter alternative

since the “OCl” limit is found independently of time and potential of overall oxide growth and so

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is hardly likely to be the result of an arbitrary steady-state situation between rate of formation of

“OC I ” species and rate of its conversion to material in an higher oxidation state.

(A) (B)

Fig. 34 Model of the relation between the Pt surface and the thin and thick films on it.

11.X Development of Thick Film Oxide States at Pt

When Pt electrodes are subjected to anodic polarization above cu. 1.9 V, RHE for

extended periods of time, substantially thick films of oxide, up to 30 equivalent monolayers of

“PtO”, can be formed as was shown in several important papers by Shibata and coworkers

] 150.15 1,174] and others [ 1481, and referred to briefly above.

In cyclic voltammetry experiments, the thick film material is reduced at substantially less

positive potentials than the quasi-2-d state from which it is usually well resolved [148,174].

Shibata referred to the latter state as the “a” oxide and the former as “p”, although that state is

itself resolvable into 3 stages in reduction, two of which arise over the H UPD potential range

[ 14X,30]. The thick film state seems to correspond tn a real stoichiometric compound, hydrous

PtO,, in which Pt is in the +IV oxidation state, as indicated by ESCA experiments

[ 186,392,193,1943.

The thick film state can also be grown in a controlled way by repeated potential cycling to cu.

2.0 V, RHE, or by restricted range cycling, as found by Burke [ 1801; similar procedures at Ru or

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402 B.E. Conway

Ir also develop thick films of RuOZ or IrO, [ 1491. At Au, visible oxide films can be generated by

strong anodic polarization [63] and with the aid of superimposed a.c.

The relation between the behavior of the so-called a and p states of the anodically formed oxide

film studied by Shibata [ 1741 and their optical behavior was examined ellipsometrically by

Gottesfeld, Yaniv and Laser [ 1981. Fig. 35 shows the relation between the ellipsometric

parameters A and w for the two states of the oxide, designated cx and p (p, the thicker film state),

over a complete cycle of formation and reduction.

I I

3 c Pt, IN H,SO, a REDUCTION ‘1

65O, 546.1

n’m2

0 CYCLI&.\‘\/

,’-1 1

0 -10 -20 -30

CHANGE OF Ll”

Fig. 35 Relation between the ellipsometric parameters A and w for the two states, a, p. ofoxide formed at Pt (from Gottesfeld et al., ref. 198), over a complete cycle offormation and reduction.

Formation and growth were conducted at 1.65 V while film a was reduced at 0.5 V and J3at

0.0 V, RHE. Gottesfeld et al. [ 1981 noted the important fact, confirmed later by the cyclic

voltammetry experiments of Birss [197] and of Tremiliosi Filho, Jerkiewicz and Conway [ 1481,

that the ellipsometric behavior indicated that the P-film behaves as though it is completelyindependent of the (underlying, cf. ref. 179) quasi-2-d film. Its earlier, well resolved, reduction

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caused almost an exact reversal of the optical changes observed during its formation at the initial

free metal surface in the fist stages of oxide film formation. They also noted that formation of

the p state of the oxide was accompanied by substantial roughening of the surface (cf. refs. 223-

232) which caused a decrease of both A and w compared with the respective values for the base

Pt surface. It was found that the roughened surface could be “electrochemically annealed” back

to its initial smooth state [148,203,204] by cycling between 0.0 and 1.5 V.

The optical behavior of the thick oxide films that can be grown at Pt anodes polarized at high

potentials (>1.8 V, RHE) for long times was examined by Gottesfeld et al. [205] in a more recent

paper. Such films are in the p state of Shibata and Sumino [200] although initially the

“ellipsometric spectrum” of the quasi-2-d state can be observed [206].

It is interesting, in relation to the cyclic voltammetry results [148,197], which show

independence of formation and reduction of the quasi-2-d state from that of the thick film

material. that there is a clear corresponding difference of the optical behavior of these states,

shown by the ellipsometry results. Also the “PtO” layer exhibits [205] strong optical absorption

over the whole domain of h scanned, bringing about a lowering of ~JI. In contrast, the subsequent

growth of the thicker oxide (represented as “PtO,“, in oxidation state +IV) is indicated by

opposite (positive) shifts of w, typical for a transparent or only weakly absorbing film with

nfihn>netecuolyte203] (cf. also refs. 70,112 and 175).

Fig. 36 shows the spectrum for q for the (“PtO”) oxide layer formed at 1.6 V. The behavior

corresponds to n-2, K-2.5 for a d value of 0.8 nm, corresponding to a “pseudo-metal” type of

oxide film state. For the fiim formed at 2.1 V, the optical spectrum of the resulting “PtO,” film

is quite different, being then more similar to that of ordinary thick oxide films at other metals

with the real part of Qoxide eing in the range 1.7 + 0.05 and K $ 0.1 to 0.2. The thicknesses of

these films can be up to 50 or so equivalent monolayers [ 1591.

The apparent thicknesses of the “PtO” and “PtO,” layers as a function of growth time at 2.1 V

were also determined, and indicated that the growth is not determined by an high-field transport

mechanism [ 1171 but rather by a “transpassive” breakdown of the inner “PtO” film allowing

oxidation of Pt2’ to Pt& at the outer surface of the “PtO” layer, leading to progressive

development of the “PtO,” film, seemingly as an hydrous structure.

An earlier paper that has attracted much interest in the area of optical studies of Pt oxide

formation is that of Ord and Ho [117] who adduced evidence for an high-field ionic conduction

mechanism (as for thick oxide film growth at other metals such as Ta and Ti) for anodic oxide

film growth, based on the relation between Tafel slopes for oxide growth and ellipsometric

determinations of film thickness as a function of time. These results apply, however, to

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404 B.E. Conway

development of the oxide beyond its “initial stages”.

3.0

2.5

n(r)

2.0

1.5

pai

I I I I I Id1.0

3.5

k(e)

3.0

2.5

2.0.0 L300 400 500 600 700 800

WAVELENGTH / nm

4.0 I I I I I I I I I

pa-j

1.0

3.5

nW

3.0

0.8

k(e)

06

2.5

2.0300 400 500 600 700 800

WAVELENGTH / nm

Fig. 36 Spectra of n and K of the single “PtO” (01) ilm as evaluated for (a) the Pt surfaceoxide formed at 1.6 V and (b) after 4h growth at 2.1 V (from Gottesfeld et al., ref.205).

The Tafel slopes were found to be proportional to film thickness (cf. refs. S&59); this indicates

that it is the field across the film (cf. ref. 20) that drives the growth process, at least for the fairly

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thick films studied in this work. This conclusion differs substantially from that of ref. 205 or of

Gilroy [ 1431 who developed a treatment based on a nucleation-and-growth mechanism, and of

Conway et al. [ 1571 who treated the initial stages of film extension in terms of a field effect on

surface-dipole potential due to place exchange, a mechanism also involved in oxidation from the

gas phase, as shown by Tompkins et al. 12071.

The critical point of interest is that a direct log law of oxide film growth (qO*proportional to log

t) applies in the early stages of growth and the relation is observed to remain continuously linear,

i.e. with a constant slope, below, through and beyond monolayer OH formation. This implies that

the kinetics are (pseudo-)zero order in OH (or 0). Only a relation such as dq,,/dt=k exp[Qq,,]

will integrate to a linear log function. Any kinetic equation of the form dq,Jdt=dq,, will give a

growth law in log q,, rather than log t. Thus, only a field effect on the reorientation of Pt/OH or

PtIO dipoles (place-exchange) leads to the required direct log law for oxide growth in its early

stages, i.e. for the compact OCl film.

II.9 Surface Structure Changes at Pt During Oxide Film Formation and Reduction

The induction of real area changes of Pt surfaces, and some anodic dissolution, by strong

surface oxidation, and/or by cycling into the “oxide formation region”, has been known for some

time. Much more sensitive and quantitative characterization of the early stages of this effect has

been made by Ross and co-workers [224-229; 231,232] by means of LEED at single-crystal, Pt

( I I 1) surfaces. Related UHV work by means of LEED and AES on effects of voltammetric

cycling has been reported in an important paper by Aberdam et al. [223].

The first LEED and AES studies of the reconstruction effects of oxide film formation and

reduction were made by Ross (2241, coupled with cyclic voltammetry. The latter technique

provides a sensitive electrochemical monitoring procedure for surface-structure changes due to the

well known sensitivity of the UPD H current profiles at Pt to surface lattice geometry (Clavilier).

The principal consequence of cycling into oxide formation potentials is the development of step

terraces; such step arrays can be characterized by LEED with the diffraction patterns calculable

[224].

Similar behavior arises at Au with cycling. However, formation of an oxide film to a surface

charge of 250 pC cm ’ (corresponding to an OH monolayer) did not produce a structure change

observable in LEED but oxidation to 300 pC cmm2 id, corresponding to the monolayer OH

deposition plus some Pt(+II) compact quasi-2-d film.

In the LEED work of Aberdam et al. 12231, ormation of disordered monatomic steps were

indicated although cycling up to 1.15 V (RHE) in aq. HClO, did not disorder the most stable,

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406 B.E. Conway

(1 1 ), surface of Pt. Quantitative analysis allowed a step-height distribution (Fig. 37) to be

derived; it is continuous in terrace widths.

Fig. 37 Step height (d) distribution derived from LEED observations of a Pt surfaceperturbed by oxide formation and reduction cycling. (From Aberdam et al.,ref. 223). Plots of q(P-hx-ky) vs (Q-hx-ky)* for the 00 and 01 beams.

Randomly distributed monatomic steps are separated by at least 11-12 compact terrace rows.

Surfaces cycled 30 times were not much more perturbed than those cycled only five times. In

H,SO,, the perturbation of Pt( 111) is more severe: the terrace-width distribution is much

narrower, down to a mean of 7.5+1 rows. Heating in vucuo at 250-300°C restores long-range

order on Pt( 111). These latter results confirm again the important role of HSO, ion adsorption in

the surface electrochemistry of Pt.

Recent STM work by Itaya et al [229,230] directly demonstrates the developed disorganization

of Pt surfaces that have undergone surface-oxide film formation. Major differences can be seen

(Fig. 38) between cycling Pt to 0.9V and 1.W. At Au, surface reconstruction can be seen under

the STM already in the “double-layer” region but this can be associated with assistance from

anion chemisorption prior to oxide film formation.

An important and general conclusion reached in the papers of Ross and of Aberdam is that the

perturbation of Pt surfaces due to oxide film formation and reduction originates from the basic

process of place-exchange which accounts for film growth beyond the monolayer stage, unlike the

situation for 2-d UPD of H or metal adatoms.

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I , I I , , , I

50 100

Ylnm

1nm

nm

Fig. 3X STM pictures of a Pt surface cycled between 0 and 0.9 V (RHE) and 0 and 1.W.

(From Itaya et al., refs. 229, 230).

Part III: COMPARATIVE BEHAVIOR TN SURFACE OXIDE FORMATION AND

REDUCTION AT Au

III. 1 Introduction

Based on a series of papers on the application and interpretation of results from use of the linear

sweep voltammetry method [55,56,57,76,77,79] at Pt, Au, Ru and Ir, a general mechanism of the

early stages of anodic oxide film formation at these metals can be developed. In this Part III of

the review, we shall base the material substantially on some papers of Hamelin et al. and our own

recent papers in co-authorship with her mainly on surface oxidation of Au. For the purpose of

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III.2 Time Evolution of the Initial Stages of Surface Oxide Formation at Au

Complementarily to the work on Pt, much work has also been reported on the processes of

surface oxide film formation and reduction at Au, including single-crystal surfaces. At Au, it

seems that the place-exchange process is more facile than at Pt (due to the lower lattice energy)

but a reversible 2-d state can be resolved at Au at room temperature (Fig. 15). However, anion

adsorption effects seem to be stronger than at Pt. Also, at Au, the time-evolution of various

stages of the oxide can be followed more sensitively than at Pt.

By means of controlled-potential anodic oxide film growth experiments, the very initial stages

of the process can be followed down to 10 to 50 us time scales with development of OH or 0

coverages down to e-OS%. The use of rapid, digital recording instrumentation in recent years

has improved the facility and accuracy with which such measurements can be made.

The procedure usually involves the determination of sub-monolayer coverages by OH and 0

species by their current/time response in a cathodic reduction sweep, following potentiodynamic,

or preferably potentiostatic, formation of the oxide to, or at, respectively, a definite potential in

the potential range within which the oxide film can be formed and/or grown. In such

experiments, the state(s) of the oxide film formed anodically during a potential holding period, b,

at a potential. E,, are distinguished cathodically in their reduction. Usually, in such reduction

sweeps initiated after polarization for tr, s at E,. more than one stage of reduction of the film can

be resolved depending on conditions. It is presumed that such resolution represents states in

which the oxide film was developed initially in response to the anodic polarization at I?,,;

however, it must be realized, as pointed out in Sections I.15 and 11.3, hat in a linear potential

sweep, peaks distinguishably observable in the reduction sweep may represent stages in the

reduction that differ for kinetic [56,57] rather than, or as well as, thermodynamic reasons.

Usually it is possible to distinguish the latter two factors through experiments over a wide range

of sweep-rates, down to low values, and from the separation of current peaks on the potential

scale, i.e. in terms for the Gibbs energies for the reduction of the supposed distinguished states of

the oxide. This is certainly the case for quasi-2-d and the thicker 3-d films that can be formed at

Pt (see Section II.4 and refs. 159,174) in aqueous media, where there are also characteristic

differences of optical properties [ 198,203,204] of the states.

Figures 39a and b illustrate the kinetic-separation effect, i.e. how, for different conditions of

sweep-rate, the sequential processes in surface oxidation of Au can be resolved essentially by

increase of the values of the s/k parameter [76,77] so that the initial, effectively reversible

electrochemisorption step is revealed at Au, as was demonstrated also in the work at Pt [55] using

high sweep-rates but at low temperature (213 K) in order to slow down the totally irreversible,

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410 B.E. Conway

post-electrochemical step of place-exchange which drains the otherwise reversibly electro-active

OH species in a 2-d array on the metal surface lattice.

RTO

Au POLYCRYSTAL

Fig. 39

I I I0.5 1.5 E/V

Cyclic voltammograms for Au oxide formation in 1O-3M aq. HClO, ata) 20 mV s-r and b) 20 V se1n 1.0 M HCIO,. Distinguishable stages in oxide

formation and reduction are identified by lettering.

In the case of polycrystalline Au in acid media, Angerstein-Kozlowska, Conway, Tellefsen and

Barnett [ 1391 studied the time development of stages of oxide film formation at Au down to 50

ps time scales and 0.25% coverages by combined potentiostatic and potentiodynamic

measurements. The general potential/time conditioning and measurement program was as shown

below:

Time

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Electrochemical Oxide film Formation 411

The initial 500 V s-’ sweep (stage 2 in the diagram above) ensures that very little surface oxide is

formed (i.e. during the sweep to the Er, potential) before the E&, rowth potential is attained.

The time evolution, referred to above, concerns the transformations that arise in the

metal/solution interphase as the oxide film develops in competition with pre-adsorption of anions

(see Section I.1 1 and Figs. iSa,b) coupled with OH or 0 place-exchange ‘with atoms of the initial

metal surface; the latter process eventually leads to desorption of the anions initially present and

ultimate development of a quasi-3-d or 3-d thicker oxide film.

Figure 40 shows a sequence of cathodic i vs E profiles for 298K taken at 50 V s-‘, following

application of the potential-time programme shown above.

Fig. 40 Sequence of cathodic i vs V profiles for reduction of an oxide film formed at Au at1.34 V (RHE) for a range of holding times t,, (from ref. 139).

The important experimental variable for the set of curves in Fig. 40 is the time, t,,, for which the

potential E,, is held constant. The example illustrated is for lZ,,=1.34 V, RHE. The holding

times, h, are in the range 2x10‘% to 5s commencing with t,=O.

The E, values were chosen to be just in the OAl region (cf. Fig. 39); for zero or small (0.005

s) t,, values; the cathodic sweeps reveal, at first, an OCl peak (cf. Fig. 39) at cu. 1.30 V; its

charge initially increases somewhat with time due to anion desorption. After 0.005 s, the OCl

charge then decreases and, over the same time interval, the OC2 and OC3 peaks begin to appear

(Fig. 39). With decreasing Q for OCl there is a shift of peak potential to somewhat less positive

values.

After ~=5 s, the QCl peak has disappeared while the state associated with the OC2 species

increases in coverage initially in a logarithmic way with time but eventually attains a limiting

coverage (shown in Figs. 42 and 39); the state associated with the OC3 peak continues to grow

in a logarithmic way with time (see below), as at Pt. This time-dependent behavior of these OC2

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412 B.E. Conway

and OC3 peaks, which arise at much less positive potentials (1.05 and 0.90 V, RHE, respectively)

than that for OCl, is what is of special interest here in relation to the transformation processes

that are taking place, in time, on the Au surface.

For certain times of holding, the i vs E profiles of Fig. 40, and similar families of plots for

other potentials, cross at an almost singular isopotential point which suggests that the process that

corresponds to a decrease of the charge in the OCl peak is intimately connected with the

processes that lead to increases in the coverages of the oxide species reducible in the OC2 and

OC3 states. Fig. 40 and similar families of plots obtained for other E,, potentials, provide a time-

resolved record of the transformation and growth of the initial states of surface oxide formation at

Au down to very low coverages.

It is informative to plot the family of i vs E profiles of Fig. 40 for increasing t,, as a quasi-3-d

diagram with t, plotted along the z-axis as illustrated in Fig. 41; this representation clearly shows

the transformation, in time, of species corresponding to the OCl peak to the other two states,

OC2 and

Fig. 41

OC3, reducible only at less positive potentials. oc2

oc3PA

Sequence of the i vs profiles of Fig. 40 plotted on a “3-dimensional” diagram to

show the time evolution of reduction-current profiles for the distinguishable oxidestates developed upon potential holding at 1.34 V. (From ref. 139).

In Figs. 40 and 41, the t,, value (for $=1.340 V) is insufficient to lead to the eventual saturation

level of the OC2 state coverage that is illustrated in Fig. 39b, which clearly shows how the OC3

oxide species continues to grow with increasing E,, while the OC2 coverage remains independent

of E , from 1.4 V onward or of time of anodization. The OC 1 species is already undetectable

beyond cu. 1.37 V, i.e. in the time-scale of the data plotted in Fig. 40; it has become completely

transformed to the OC2 and OC3 states of the oxide film.

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The significance of the OC2 peak still remains to be elucidated. It reaches a saturation level of

CU.50 pC cm-* or equivalently 8,,- 0.25, and is only resolved in CIO; solutions, not SO,” or

HSO,, so its absence in the latter solutions must be associated with strong anion re-adsorption as

free metal s ites are generated upon film reduction. HSO; added after film formation in dilute

HClO, also leads to absence of the resolved OC2 state, so this is not some state generated

specifically &rin~ the oxide film formation in sulfate solutions.

It is to be noted that the time-scale of the oxide film formation and transformation processes s

enormously different at lower temperatures, e.g. 274 K, where all processes are substantially

slowed down, indicating that a substantial activation energy is involved.

The time-dependence of total oxide reduction charge, Q,,,,, and of its OCl component, Qor,,

plotted as f(log t,J in Fig. 42, shows that for E,=1.26 and 1.27, and probably for 1.28 V, the OCl

peak charge remains constant with t, over 3 decadesof

01 M HCIO,

s = 5.10-2 V s-’

POTENTIAL, E,/V

time.

-ml

Fig. 42 Charges (Q,,,) for reduction of the oxide film at Au, formed at 1.26 V and beyond,

and its components (Qoc, and QOC2) s a f(log t& T=298K. (From ref. 139).

This virtual constancy of Qocl values with time signifies, that at these sufficiently small anodic

potentials, the state of OH species (corresponding to the OCl peak in reduction) is stable and

almost at equilibrium, corresponding to a simple charge-transfer reaction involving reversible

deposition/reduction at the surface. In this way, the process is similar to that for the initial stages

of OH deposition at Pt 1551. Only beyond 1.28 V (at this experimental temperature of 298 K)

does the OCI state of electrosorbed OH begin to be transformed to states of greater stability (cf.Fig. 41), as manifested by their smaller positive reduction potentials. Of course, at higher

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414 B.E. Conway

temperatures, the transformation becomes more facile and is detectable at smaller f values and

lower potentials.

On the basis of the mechanism of the anion adsorption/desorption process in surface oxidation

summarized earlier, it is useful to explore the extent to which changes of coverage by the species

corresponding to the OCl peak (Reaction 2 in Scheme 2; see later), are balanced by the

increasing coverage by the OC3 and OC2 species. Such a balance is suggested by the

observation of the isopotential points that arise in families of i vs E profiles for reduction of

states of the oxide formed after various intervals of time, as exemplified in Fig. 40.

An application of this charge-balance evaluation is illustrated in Figs. 41,42 for l&=1.340 V for

the same holding potential. Similar diagrams of charge balance can be drawn for oxide growth at

other potentials. As can be seen in Figs. 40.41, an initial slow growth of the charges in the OC2

and OC3 peaks is accompanied by only a small decrease of the OCl peak charge. This is

followed by an increase of charge that is logarithmic in tt, for both the OC2 and OC3 peaks. The

OCl peak decreases also logarithmically in t,, over the same range of time.

The slope dQ/dlog ti, is, in the case of the OC2 peak, independent of the potential of growth.

Only the time on the log b scale, at which the logarithmic growth begins, becomes smaller with

increasing E,. A point of substantial interest is that the growth of the OC2 state reaches a

sumration value which seems to depend very little, if at all, on the potential of prior film growth.

The slope dQ/dlog t,, for the OCl peak does not seem to depend on E, nor on time, for anodic-

sweep end-potentials positive enough for the reversible state OAl/OCl to be fully developed.

When the growth potential is such that only a fraction of the full charge for this peak can be

transferred, the decrease of charge starts at longer times, but it joins eventually the common Q vs

log f line.

The ratio of charges during the conversion of OH, deposited in between the initially

chemisorbed anions (Reaction 1 in Scheme 2), into the MOH and MO states can give information

concerning the nature of the transformation.

III.3 Physical Basis of the Transformation Process

It is evident that OH species can be initially electrodeposited reversibly in 2-d arrays involving

a “preadsorbed” anion sub-lattice [Reaction(3) in Scheme 21, depending on the E,, potential and

temperature, but the species undergoes a time-dependent, “post-electrochemical” process which is

identified as place-exchange, as discussed earlier. The process is coupled, as a concerted

reaction, with anion desorption in Reaction (4) (Scheme 2). Compared with Reaction (3), it is

usually highly irreversible. In the place-exchanged state, the OH or 0 species (now represented

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as OHM or OM) can only be reduced at less positive potentials than those for the OCl region, so

“hysteresis” in the cyclic-voltammetry curves arises.

xM+A~.zH,O+M,A-.sH20 (Electrostatic adsorption of anion) (1)

M,A-.sH,O~,A”-~‘-.~H,O+&- (Chemisorption of anion) (2)Then, for &I:

M,A.yH,O~,A.OH~“Y)-.(y-I)H,O+H++~-. Discharge of OH from

hydration water in(3)

preadsorbed anion sub-lattice

M,A.yOH”-Y’-+yM+yMOH+~+A-+[y( I-y)-l]e- (4)

+zH,O+e-

A-.zH,O

(in solution)

+xH,O

Completion of OH discharge, change of H-bond environment ofanion, with anion desorption coupled with place exchange

x(MOH+H++e-)

MOH+ MO + H++e

J L

“OHM” “OM” Place-exchange

(5)

(6)

(7)

Scheme 2

The place-exchange process continues in time and is enhanced in rate with increasing e lectrode

potential, E ,, or corresponding field (see below), and with temperature. The extent of formation

of place-exchanged species, characterized by the charge under the OC3 peak, measures the degree

to which production of a quasi-3-d phase oxide has taken place; eventually, with sufficient time

and/or at higher potentials and elevated temperatures, a true 3-d phase oxide can be formed at Au

and other noble metals (Part tI).

For longer t, and higher &, Qo,, is much larger than Q,,, (Fig. 42) and corresponds (at Au) to

the main continuing oxide growth process but, as Fig. 42 shows, at %=1.34 V and for t,, up to 5

s, the coverages by the OC2 and OC3 states increase comparably, but eventually the OC2 state

reaches its saturation limit of coverage substantially below a monolayer, as mentioned earlier. At

the moment, it is unclear what is the significance of the attainment of saturation of the OC2 state

which develops befke the OC3 becomes substantially formed (Figs. 40,41). OC2 is not

separately resolvable if the electrolyte is moderately strong H,SO, (-OSM), so there is an HSO;

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416 B.E. Conway

anion re-adsorption during film reduction which then blocks the resolution of the OC2 state.

These results provide the basis of the evidence that the initial stages of oxide film formation at

noble metals are far from simple, and various coupled adsorption/desorption and reconstruction

processes take place prior to (bulk) “oxide film formation ‘I in the conventional sense. Anion

“pre-adsorption” is especially important (see below). At Au, these processes can be resolved and

followed rather more sensitively than at Pt while, at Ru and Ir, multi-layer film growth, being

more facile, interferes with the study of the initial stages. In fact, e.g. at Ru, the initial stages as

revealed in a cyclic-voltammogram, can only be resolved in the first one or two cycles on a fresh

electrode surface so that care must be exercised in such an experiment not to involve the

formation of the multilayer oxide film that exhibits reversible “redox” behavior like that also at Ir

[ 1461, as mentioned earlier. It is interesting that thick oxide films can also be formed at Pt, Pd

and Rh by cycling or ac. modulation procedures but the films never show the reversible redox

oxidation/reduction behavior observed at Ru or Ir, or even Co in alkaline solution.

III.4 Comparative Cyclic Voltammetry Behavior of Principal Index Planes of Au in Aq.

HCIO,, and the Role of Anion Adsorption

As in most aspects of the surface science of metals, more informative and specific results are

obtained by use of well defined surfaces of single-crystals. This is the case with surface oxide

formation at Au, notwithstanding the reconstructions that can occur [223-2321. In fact, surface

index-specific behavior is observed, analogous to that for H UPD at Pt surfaces.

Before reviewing details of the oxidation behavior of the principal index planes of Au and the

specific adsorption of anions thereon, it will be useful to summarize the main physical

characteristics of these surfaces as given in Table 111.1.

The i vs E profiles for the (1 11 , (100) and (110) planes of Au in 0.0 1 M HClO, are compared

in Fig. 43. Even without detailed analysis of the profiles, rather large differences in the oxidation

side of the profile are easily seen in the following regions:

i) at the onset of oxidation, due to the different extents of competitive adsorption by anions;

ii) an the region where reversible deposition and reduction of OH arises (peaks OAl/OCl) in

between specifically adsorbed anions [ 1051;

iii) in the shape and number of the RTO (place-exchange and anion replacement, turn over)

peaks).

Essentially there are no differences amongst [223,234] the crystal planes in the main reduction

peak OC3, when reduction occurs from a monolayer or more of OH. Differences are, however,

again observed in the OC2 peaks where reduction of OH takes place from (in HClO,) the anion-

8/7/2019 Con Way 95



free surface after most of the quasi-2-d oxide (OC3) has been reductively removed.

Table III.1 Characteristics of the low index single-crystal planes’ of Au

AtomicUnit cell diameter Symmetry Number of

No. of Area a (A) and TLK atoms per QdS”Plane atoms (A’) notation cm2 (pC cm-2)

(Ill) 1 a2\/3/2=7. 19 three-fold 1.39x1o’5 222

(100) 1 a2 = 8.34 2.887 four-fold 1.20x10’s 192

(110) 2

(1)

a 2=11.79J two-fold 1.70xlo’5 272

2(111)-(111)(0.85~10’7 (136)

PolycrystallineAu

1.25x10’s+ 200+

*For these and other structures, see J.F. Nicholas, An Atlas of Models of Crystal &faces,Gordon and Breach, New York (1965).

‘Calculated from the charge Q required to form a monolayer of Au0 [213].

Even this superficial comparison of the i vs E profiles in Fig. 43 suggests that a different

configuration of adsorption of anions, relative to the lattice geometry of the surface, may be the

cause of the differences observed. Then, in the so-called “double-layer” potential region,

differences should also be seen in anion adsorption. Jndeed, as is seen in Fig. 43, this is the case

and the i vs E profiles over this region vary characteristically in shape and charge passed with the

symmetry of the planes. However, on each of these planes [excluding the (100) plane], the

charge passed surpasses. although to various extents, that expected (qd,) for electrostatic chargingof the double-layer (D) over this potential region. The passage of this excess charge leads to the

important conclusion that Au is not ideally polarizable in this potential range because there is

specific adsorption of the anions accompanied by partial or full charge transfer. The peaks

observed in this potential region are denoted as Dl,D2, etc. for the anodic profile and DC for the

cathodic profile for the anion adsorption curves.

The specific stages of oxidation of the single-crystal planes, mentioned above, will now be

described in more detail.

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418B.E. Conway

Fig. 43.

RTO

aa 2-lI 2-2

c

RTO

M OA 2*

Comparison of cyclic voltammograms for Au( 11 l), (100) and (110) surfaces in

0.01 M HClO,. (From Angerstein-Kozlowska et al., ref. 223.)

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III.5 Sequential Stages of the Surface Oxidation Reaction at Single-Crystal

Surfaces of Au

i) Peaks for reversible processes and role of anion co-adsorption. It was shown in Part I

(Sections 10 and 11, and Fig. 15) how the initial stages of OH deposition in oxide film formation

take place on a surface that already bears a sub-monolayer of chemisorbed anions.

This process of deposition of OH in between anions already adsorbed on the surface, is seen as

a reversible current peak or peaks at potentials less positive than that for which the main

oxidation of the surface commences. It has different character on the various single-crystal

planes studied.

On the (100) plane in HClO,, one rather broad peak (Fig. 43), the so-called “preoxidation”

region OAl, is observed, the shape of which indicates significant repulsive interactions between

the anions and OH on the surface suggesting [215] incomplete charge transfer on adsorption

rather than adsorption on energetically different sites in between an overlay-lattice of adsorbed

anions on the surface. The latter situation arises on the (111) plane where the two “reversible”

peaks, rather close together (OAllOCl and OA2/OC2, Fig. 43) seen in HClO, do not indicate

(from their widths) any effect of repulsive interactions but suggest rather, that OH is deposited in

two different energy states on the surface, probably associated with sequential stages of lattice

filling and coupled anion desorption.

The lack of excess charge q~, which could be connected with anion adsorption on the (100)

plane in dilute HClO,, also indicates that the ClO,- anion retains its charge upon adsorption on

this plane (Table 111.2).

An excess charge, q,, is, however, seen for this plane in H,SO, which has strongly adsorbable

anions, and in both acids on the other planes.

When strongly chemisorbed anions of H,SO, (normally HSO,-) are present, the preoxidation

peak on the (100) plane is completely blocked as can be seen in Fig. 43 and Table 111.2,while,on the (111) plane, the peaks are only shifted to more positive potentials [%I, and deposition into

a third class of energy sites is seen in a third small peak before replacement of the anions by OH

in an RTO peak starts [56].

Around the stereographic triangle, between the apices corresponding to the principal index

planes, stepped surfaces arise. Much interest has arisen in surface electrochemistry on special

properties of such surfaces for electrosorption and electrocatalysis. Two important papers by

Adzic et al. [242,243] have addressed oxide formation on Au stepped surfaces [243] andadsorption of H and HSO,‘ or SO,* ions on such surfaces at Pt. A pronounced structural

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420 B.E. Conway

sensitivity for both oxide film formation and anion adsorption is found at Au, with correlation to

the double-layer capacitance. For some of the surfaces, vicinal to (11 l), a substantial separation

between the oxidation at the step edges and the adjacent terraces is seen.

Table III.2

Solution 0.01 M HCIO, 0.01 M H,SO, 0.1 mH,SO,

Adsorbing

speciesPlane

(111)Peak(s)OA3

(100)

Peak(s)(100)Peak(s)

ao,-

4a(PC cm-“)6-10Dl-2

0

D5-6DlD2

OH- HSO,-SO,‘- OH- HS0.,-,S0,2-

.%H q.4 %H %(PC cm-2) (pC cm-2) (PC cm-2) (pC cm-‘)

35 32 63 43

OAl-OA2 Dl-3 OAl-OA2-OA3 Dl-3

20 13 0

“preoxidation” D “preoxidation” n/a16-18 n/a n/a 22

Dl D2

%H(pC cm-7

63OA I-OA2-

n/a0

n/a signifies data not available; k,, the charge for OH deposition.* aA=q-q,, where qdl was calculated for all the planes in HClO, and for the (111) and (100) planes in H,SO,

solution taking C,=4OpF cm-2, as was measured at the positive end of the double-layer potential region before thepotential for the main oxidation process on these planes was reached. For the (110) plane in H,SO,, C, was only ca 32 ~JFcm.2. and this value was used for calculations. Because of the pseudocapacitance associated with anion specific adsorption

with partial charge transfer, a background value of the electrostatic capacitance of the double-layer is difficult to estimateaccurately.

In the case of stepped Pt surfaces [242] from the (110) zone, HSO,. and SO,” are

strongly adsorbed at the trigonal sites at the steps and the H adsorption is correlated with

the sulfuric anion adsorption, as at the (111) plane.

On the (100) plane, no obvious reversible oxidation peak arises although, as shown in

Table 111.3,a peak D2 having the same general characteristics as the preoxidation region

on Au (100) can be observed in the double-layer potential region in HClO, (Fig. 43, Table

111.2). This process is completely blocked in H,SO, (Table 111.2). It is interesting to notethat for the (110) plane in HClO, solutions, the charge, q,,, in peak D2 is 3 times the

charge q, for anion adsorption in peak Dl.

ii) Peaks for irreversible processes. The process of replacement of the adsorbed anions

by OH, to which the RTO peaks are assigned, requires different energies on the three

planes. In HCIO, it is easiest on the (100) plane where both sharp RTO peaks, OA2- 1

and OA2-2, arise at less positive potentials than the RTO peaks on the other planes (see

Table 111.3).

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Table III.3 E, (V) for Au low index planes in 0.01 M HClO,

421

Potentialregion

Double-layer Oxide formation

Peakreversible RTO

Dl D2 D3 OAl OA2 OA2-1 OA2-2 OA3 OA4

Plane(111)( 100)( 110)

0.840 1.16 1.22 1.275 1.340 1.5600.81 0.94 1.16-7 - 1.28 1.32 1.40 1.59

0.73 0.92-3 1.31- 1.36- 1.44 1.541.32 1.37

The fact that these peaks are so narrow indicates that the removal of an anion, little

discharged on (100) causes the local potential suddenly to change to a more positive value

with the result that a substantial number of OH’s are deposited on the area previously

blocked by the anion, with a sharp increase of current. This large area (evidently blocked

by the anion) and the facile replacement of the anions is in accordance with the

indications from the lack of excess charge in the double-layer potential region, that the

anions on the (100) plane are very little, or not at all, discharged on adsorption.

The situation seems to be different for the (111) plane where replacement of the anions

is most difficult (Table III.3). The RTO peak here is OA3 at 1.34 V, i.e. the energy

corresponding to completion of the third sublattice of OH is required for replacement of

anions on the surface by OH. This is probably because the trigonal geometry of HSO; or

ClO, is geometrically commensurate with the 3-fold symmetry of the (111) surface. On

the other planes, the energy corresponding to formation of the second sublattice is

sufficient for the OH/anion replacement reaction to take place.

The passage of the second electron in the OA4 peaks occurs in HClO, solutions at

similar potentials on all three planes. The peak is rather broad on the (100) and (110)

planes, and only on the (111) is it well pronounced. On the other planes, for which the

symmetry of the surface arrangement of atoms is different from the symmetry of the

anions. the desorption of anions is completed much earlier in the RTO peaks (OA2) before

even the OA3 peak is reached. Strong adsorption of HSO,- anions is also indicated by the

much larger excess charge over the double-layer potential region than for other planes,

which amounts to cu. 43 uC cm-* for 0.1 M H,SO, (Table 111.2).

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422 B.E. Conway

III.6 Concentration (Activity) Dependence of the E$ Values as a Basis for Evaluation

of Anion Adsorption with Charge Transfer

i) HCIO, solutions. By analysing peak potential data in the anion pre-adsorption

potential range as a function of anion concentration (activity) in terms of a Nemst

equation for the chemisorption process with charge transfer: A-+M f M.A”-S’m+Se(M),

information can be obtained on the value of 6, the extent of charge transfer accompanying

anion chemisorption (cf. Lorenz and Salit, ref. 60). A similar procedure can be applied to

the analysis of OH electrosorption: H,O + M + M.OHoY’- + y (Scheme 2, p.98).

When such measurements are made as a function of log [acid concentration] using a

Pt/H, reference electrode, allowance must be made, of course, for the H+-concentration

dependence of the RHE, giving a “constant potential electrode” (cpe) scale. Fig. 44 shows

such results (on the RHE scale) for E, values of the reversible processes at the various Au

principal index planes, while Fig. 45 shows similar results for the non-reversible process

peaks.

PEAK F’LANE

Fig. 44 Peak potentials for various reversible processes at the Au(II1) surface as afunction of log concentration of acid.

In all the cases a linear relation between E, and log c is observed. The resulting slopes

generally fall into three categories: they are equal to zero or are positive for the oxide

formation, OA, peaks and they are negative for the (anion) D peaks in the double-layer

potential region, (Fig. 46).

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Fig. 45

Fig. 46

1.6 a(lll)

-x 1 I-

1.4

-0-o

brtoo,x-

+-.-

E&l

OA4

OA3

OA4

oA3

oA2-2OA2-I

As in Fig. 44 but for the non-reversible processes.

DI

Cyclic voltammograms at Au for the “preoxidation” anion adsorption peaks(D) in the double-layer region at potentials less than those forcommencement of OH deposition. Sweep rate 20 mV s-l, 298 K.

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424 B.E. Conway

An exception is the D2 peak of the (110) plane which has a larger positive slope than

that of the “preoxidation” OAl peak of the (100) plane and the OAl/OA2 reversible peaks

of the (111) plane. This indicates again that a reaction involving OH deposition (as in the

OAI peaks) has to be assigned to this peak.

Interpretation of the Nernst slopes, dEr/d log c, must take into account the variation of

the HJH+ reference electrode potential with acid concentration (activity), noted above.

Therefore, in the text and tables, all the values of the slopes B,,=dEr/d log c are vs cpe.

These values of the slopes on the cpe scale, from Figs. 44 and 45, are summarized in

Table 111.4.

Table III.4

Potentialregion

PeakPlane

(111)(100)

(110)

Double-layer

D Reversible

-6OklO (DC) +l lO(OA1)+105(OA2)+105

(“preoxidation”)-60 (Dl) +150 (D2)

B,,,=dEr,/d log c (mV)

Oxide formation

RTOOA2 OA3 OA4

- +95 +60 and more +ve+60,60 +60

(90)+60 +60

For the double-layer peaks, the dependence of E , on log c,,,,, can be best established

for the (I 10) plane for which the pseudocapacitance peaks Dl and D2 are unobscured by

any overlay lattice reconstruction peaks and well separated for c 2 0.01 M (Fig. 46). The

slope for Dl on the cpe scale is -60 mV and for D2 +150 mV. On the (111) plane the

anodic pseudocapacitance peak in the double-layer region changes its shape after a certain

amount of adsorbed ClO,- anions is discharged. Hence the resulting peak that is clearlyseen is not a pseudocapacitance peak but arises because of rearrangement. No

rearrangement is observed in the cathodic profile, therefore in Fig. 46 the change of the E,

of the cathodic pseudocapacitance peak, DC, with log ~u,-,,,~s shown. The resulting slope

is -6O+lO mV on the cpe scale.

On the (100) plane, as was said before, no pseudocapacitance peaks in the double-layer

region are observed. The small peaks D (Fig. 46) which appear before the preoxidation

peak are not due to anion discharge as both have a positive slope on the cpe scale; themore positive one with a slope of 50 mV follows nearly exactly the pH change; the more

8/7/2019 Con Way 95


Electrochemical Oxide film Formation 425

negative has a slope of cu. 30 mV. Hence, in both cases, OH must be involved in these

reactions.

With change of c,,,o,, the pH of the solution will change simultaneously with the

concentration of anions. From the few reactions which could possibly occur in the

investigated potential region, only the formation or reduction of surface oxide according to

M+H,O - MOH+H++e is pH-dependent in the same way as the RHE: viz,

dEdd log c,+=2.3RT/F. Hence, for a reversible-process peak produced by the above

reaction no shift of E , with change of pH should be observed when E, is measured on the

RHE scale (or B=+60 mV on the cpe scale).

None of the reversible process peaks (Table I11.4) shows, however, such a dependence,

but, the reduction from the anion-free surface in the OC2 peak at the (110) plane, on

which the anion adsorption is weak, does have the same dependence on the acid

concentration as the RHE. This is also the case on the (110) plane for the formation of

oxide in all the peaks for processes that are not reversible (Fig. 49, and on the (100)

plane in dilute solutions of HCIO,.

Evidently, on these planes, weakly adsorbed anions, although present, do not change the

energetics of these processes, as was also observed for polycrystalline Pt where quasi-

equilibrium was postulated in the oxide formation reaction with a post-electrochemical

turnover process as the current-determining step. Here the accessible space, as measured

by qml, will be influenced by the anions. This is not the case for the (111) plane on

which anions are adsorbed more strongly. and are hence displaced only with more

difficulty at higher OH/O coverages.

A positive slope, larger than that resulting from the pH-dependence of the RHE,

indicates that the formation of the surface oxide in the presence of adsorbed anions is

accompanied by a reaction, the product of which includes the desorbed anions, or that less

electrons are passed than H’s produced.

In the second case, deposition of OH without full charge transfer, probably

coordinatively with the anion, in a surface lattice process such as

MA”-6)-+M+HzOJMA(‘-6’-.MOH(1-Y)-+H++ye (44)

must be postulated. Here the electrosorbed OH species are written with partial charge

(I-$-. This process gives a dEJd log c slope of 2.3 RTIyF against a cpe. The observed

values (see Table 111.4)are cu. 102-120 mV or higher, so y must be 0.5-0.6. (cf. refs.

105,139 and 216).

The partial retention of charge by specifically adsorbed OH should cause substantial

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426 B.E. Conway

repulsive interactions on the surface, broadening [87] the adsorption peak. This is the case

for the “preoxidation” peak for the (100) plane where it seems that 6=0 and the slope of

the relation in Fig. 46, equal to 105 mV, indicates that y is cu. 0.6.

On the (111) surface in HClO,, the OAl/OA2 peaks are surprisingly narrow and thus

well resolved (Fig. 43). However, the dipolar “MOH” species to which they correspond

are deposited amongst fully or almost fully discharged ClO,- ions, so that repulsive

interactions would be expected which would tend to broaden the peaks [77]. The fact that

good resolution is actually observed leads to the suggestion that it arises for a special

reason such as H-bonding of water molecules, from which the OH’s are generated, to the

discharged ion array. The resulting discharged OH species will probably remain H-bonded

to 0 atoms of the electrosorbed ClO, particles, giving rise to an attraction situation that is

known [77] to result in narrow pseudocapacitance peaks.

On the (110) plane both mechanisms may be operating: the D2 peak, characteristics of

which are as for the “preoxidation” peak at the (100) plane, may correspond to reaction of

discharged with OH not H-bonded to the anions: the reversible peaks observed at the

beginning of oxidation, as for the (111) plane, may arise from H-bonded OH on the

surface. This would indicate that these peaks are probably an intrinsic feature of the (100)

plane.

In the first case, a competitive ion-displacement process such as

M,A”~s’.MOH”~Y)~+xH,O~( 1+x)MOH+xH++A-+(x-8+$e (45)

was postulated, where x is the number of surface atom sites occupied by specifically

adsorbed anions. This process will have an opposite dependence of its peak potential on

log c to that of the process of electrosorption of A- (eqn. 45), which should lead to what

we have called “RTO” and thus correspond to a large peak of the irreversible type.

Therefore it is not likely that such a process as (45) can already occur in the OAlpotential range.

The possible Nemst slopes associated with this process are given by

dEr,/d log c=x 2.3 RT/(x-6+y)F+2.3 RT/(x-G+y)F

assuming that OH coverage at the peak is independent of cnc,04and noting c~+=c,-,~~-=c.

The value of dEJd In c, according to eqn. (46) can be estimated for various values of x

using the 6 and y calculated from the experimentally derived slopes for the D peaks for

anion adsorption and for the peaks for reversible electrosorption of OH at the variousplanes (Fig. 43). The results of such calculations are summarized in Table 111.6.

(46)

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6 Y


Table III.6

X dEt,/d log c

427

1 0.5 3 96

6 769 71

0 0.5 1 802 723 686 649 63

0.6 1 752 693 676 649 62

It is significant that with increasing shielding of sites, i.e. for larger x, the slope dEt,/d In c

approaches 60 mV.

For a 3-site adsorption of anions with full charge-transfer (6=1) and with ~0.5 for OH

electrosorption, a slope of 95 mV arises from eqn. (46), as actually observed for the RTO peaks

of the ( 111) plane in HClO,.On all the three planes that were investigated by Kozlowska with Hamelin et al. [57], the E, for

the OA4 peaks are independent of c,,,,,~ on the RHE scale (Fig. 45); for these peaks oxidation is

therefore proceeding on a surface covered with OH and, with the exception of the (111) plane,

free of anions. For the (111) plane, in more concentrated HClO, and in H,SO,, some remaining

effect of anion adsorption on E, can be clearly seen.

In the double-layer potential region, the first reaction to be considered is a specific

adsorption of anions with charge transfer, which will be influenced only by theconcentration of anions as is obvious from eqn. (44) shown earlier, where O&i<1 is the

extent of charge transfer [60,61].

For a reversible process, the Nemst equation gives a shift of E, in the opposite direction

to that for the case of the oxide formation reaction. The slope B,=dEdd log c,. =

-2.3 RT/&F on the cpe scale i.e. it is negative and it will have an even larger numerically

negative value on the RHE scale (-2.3 RT/GnF = dEr,ld log $+)A With decrease Of G(for

6>0), the numerical values of the negative slope will increase as can be seen in Table III.7

where values of the slopes on the cpe scale are given for various 6 and n values.

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428 B.E. Conway

Obviously, for the case of adsorption with F=O, no Nemst equation for the charge transfer

adsorption can be written. The anion adsorption is then determined by the electrostatic

state of the double-layer.

Table III.7

6xnB,=dEr/d log c

(mV) On planeObservedin solution For peak Remarks

-60 (110) HClO, Dl1x1=1 -60 (111) HClO, DC

1x2=2 -30

1x1.5=1.5

2/3x2=4/3

1/2x2= 1

-40

-60

(111) WQ D3

(111) H,SQ D3

Mixture of

adsorbed HSO,-and SO,”

Comparison of the data in Tables III.6 and III.7 suggests that ClO,- anion adsorption on

the (110) plane in peak Dl takes place with full charge transfer (6=1) as well as for the

desorption peak DC for the (111) plane. The results for the (111) plane in H,SO, are

interesting insofar as, for n=lS, they could suggest a value of the ratio of coverage of

HSO,- and SO,‘- on the surface.

An increase of concentration of HClO, from 0.01 M (Fig. 47) to 0.1 and 1 M causes a

positive shift of the anodic peak potentials without a change of the overall shape of the

i us E profile, as shown in Fig. 47.

(III)

HC104

SC-

c:

5

3.-

o-05

Fig. 47 Effect of increase of ClO; anion concentration (as HClO,) on the anodicpeak potentials for Au oxidation at 298K. Sweep rate 20 mV s-l; 298 K.

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The peak potentials as a function of log ~uc,~~ re as shown earlier in Fig. 44. As

the potentials are measured on the RHE scale, any shift observed is due to the effect of

anions, not pH (cf. the discussion of this factor given earlier).

The biggest effect of concentration (c) of anions is seen, as would be expected, for the

first reversible process peaks, OAl and OA2. For these peaks, dEJd log c on the cpe

scale are respectively +I 10 and +105 mV, while for OA3 peak this derivative is only 94

mV. The implications of these high slopes have been discussed already in Sections III.3

and III.4 in terms of OH deposition within an overlay lattice of adsorbed anions with a

partial charge fraction, y, of only cu. 0.5.

The clear observation on Au( 111) of tu’o reversible surface processes (Fig. 43) indicates

the formation of two sublattices of different energy in the very initial stages of surface

oxidation.

The high value of dEr,/d log c (Fig. 44) for the large OA3 peak, points to an RTO

process in which a desorbed anion is one of the products of the reaction:

M,A”~6’~.MOH”‘Y’~+xH,0~( l+x)MOH+A-+xH++[x-&( I-y)]e-

ii) The i vs E profiles in aq. H,SO,. For H,SO,,, as can be seen from Fig. 4X where

the i vs E profiles for 0.01 and 0.1 M H,SO, are compared with the profile for Au in 0.01

M HCIO,, the onset of oxidation is shifted positively and the shift is larger for higherconcentrations of H,SO,. 100r (100)

Fig. 4X Comparison of voltammetric i vs E profiles for Au oxidation in 0.01 and0.1 M aq. H,SO, in comparison with that for 0.01 aq. HClO,.

(47)

The strongly chemisorbed HSO,- and/or SO,” (depending on pH) anions block nearly

completely the otherwise distinguishable passage of the first electron, and the RTO process

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430 B.E. Conway

takes place mainly over the potential range of the OA4 peak, where the main extent of the

oxidation of the surface (427 pC cm-‘) occurs, rather than in the OA3 peak as is the case

in HClO, solutions. The vestigial OA3 peak is seen only as a shoulder on the ascending

part of the OA4 peak. The extent of charge transfer in the OA3 peak is difficult to

evaluate; it seems to be about the same as that for each of the peaks for the first two

sublattices, as seen in Table III.%

Table III.8 Charges for designated surface oxidation processes at Au

Peak Ideal Au( 11 ), anion free Au( 111) in 0.01 M HClO,, (experimental values)

tQls Y

(hypothetical)

OAI 74 0.3(3) 37 0.17 0.5 0.34OA2 74 0.3(3)OA3 74 0.3(3) 125 0.56 l(O.5) 0.35”

222 1 oo 162 0.73 0.69OA4 222 1.00 joJ 1.35 1 1.35Total 444 2.00 462 2.08 2.04

“The charge required for desorption of anions was subtracted (see Table III.9 below).

For the deposition of OH in between anions in the overlay lattices of chemisorbed HSO,- or

SO,*-, two distinguishable energy states for OH deposition are now resolved (Fig. 48; OAl and

OA2) but there is no anodic/cathodic reversibility of these processes; instead, as can be seen in

Fig. 48 (first two curves of anodic sweep profiles), the reduction occurs at less positive potentials,

partially in the OCl’ peak indicating some desorption of anions and partially from the turned-over

state in the OC3 process peak. The ascending part of the first peak (OAl) in H,SO, is much

steeper and the two peaks, OAl and OA2, arise over a narrower potential region than for 0.01 MHClO, solution. These results suggest that, contrary to the behavior in HCIO, solution, in the

time scale of the sweep, some small but significant desorption of anions must occur before any

OH can be deposited or that attractive interactions (via H bonding) are much stronger than in

HClO, solutions. The total charge (Table 111.9)passed over the two peaks prior to development

of the RTO peak, is diminished to only cu. 46 uC cm-*,

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Table III.9 Charges for OH coverages at Au (111) in 0.1 M H,SO,

431

Peak %A %spC cm-* (hypothetical)

Y 8OH/y

OAlOA2

46 0.21 0.5 0.42

OA3 cu. 22 0.1 0.2OA4 427 1.71 l(O.5”) 1.41

Total 495 2.02 2.03

“Of the 427 pC cm-’ of the charge of the OA4 peak, 68 pC cm-’ will be used to complete thedischarge of the OHoY’- deposited in the previous peaks in the complexes with the anions. Forthis part of the charge, y=OS. For the remaining charge, 359 pC cm-*, ~1. Part of this charge(47 pC cm-*) will be used to desorb the anions and only 312 pC cm-’ will then be associated withfurther deposition of OH or 0.

III.7 Anion Cbemisorption with Charge-Transfer at Au(ll1)

Various aspects of the evidence for pre-adsorption of anions prior to onset of surface

oxide formation were given in Sections III.3 and fII.4. More directly, cyclic

voltammograms taken over the “double-layer” region at Au( 111) prior to onset of the

initial, OAl, stage of OH deposition clearly indicate a chemisorption process of adsorption

with charge transfer both in the case of ClO,- and HSO,- ions (Fig. 46). In fact, the

pseudocapacitance curves are reminiscent of UPD profiles for metal ad-atom deposition

and desorption where several sub-monolayer states are resolved at single-crystal surfaces.

A large charge, qu is passed essentially reversibly, giving rise to a group of peaks

designated D (Fig. 46) that are particularly well developed in H,SO, solutions.

qt, exceeds the charge, 36&5 pC cm-*, needed for electrostatic charging of the double-

layer over this potential interval (0.4-1.3 V), calculated on the assumption that in the

region positive to the P.Z.C., C,,-40 pC cm-*, as measured at the most positive end of the

region. The excess charges for the various solutions calculated in this way, are: for

0.01 M HClO, 6-10 pC cm-*; for 0.01 M H,SO,’ 34 pC cm-* and for 0.1 M H,SO, 47 pC

cm-*. These charges would, of course, be even larger if the electrostatic C,, were ~40 pF

cm-‘.

The origin of the extra charge cannot be linked with the UPD of OH because: (i) the

potential at which the extra charge is passed does not follow the pH change, as can beseen in Fig. 46a for HClO, solutions and in Fig. 46b for H,SO, solutions; (ii) the charge

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432 B.E. Conway

(48)

149)

increases when more adsorbable anions are present in the solution. The behavior can best

be explained on the basis, referred to earlier, that ClO,,- and SO,” (or HSO,-) ions

undergo some appreciable degree (6) of charge-transfer upon their adsorption,

xM+A-+M A(‘-6)-+ 6e-The peak potential of a region of reversible currents associated with process (48) should

follow the concentration of the anions. c,., according to a Nernst equation:

E,=E, - (2.3 RT/SnF) log c,. t

In HClO, solutions (Figs. 44,47) the potential of the cathodic peak. DC, follows eqn.

(49) for 6=1. For the H,SO, solutions, this is the case only for the D3 peak and the

descending part of the D2 peak (Fig. 46). On the RHE scale, the D3 peak shifts -100 mV

when the concentration of H,SO, is changed from 0.01 to 0.1 M. The corresponding pHchange is 1.84-0.95=0.89, which gives 0.052 V as the change in the reference electrode

potential. Therefore, on a cpe scale, dEr,/d log c,. is -0.048 V for the D3 peak. This shift

occurs in the correct direction and with the right magnitude for anion chemisorption with

6n=l. 1, i.e, with more than 1 electron transferred.

In H,SO,, the potential of the other peaks, Dl and D2 (Fig. 46), are independent of pH

and concentration of anions, indicating that almost maximum coverage by adsorbed anions

is reached, xM+A- + M,A,,, before charge-transfer occurs, viz. Mx&~~- + M,A”‘s’-+Se-and no further adsorption takes place during charge transfer. This compares well with

previous results for polycrystalline Au for which saturation coverage with anions is

reached at cu. 10e3M H,SO,. The sharpness of the Dl peak (Fig. 46) suggests a 2-d

phase transition in the ad-layer of HSO, anions in the presence of co-adsorbed water

molecules (see below).

III.8 Role of Water of Hydration of Ions in the Ad-layer at Electrode SurfacesIt is now well accepted that solvent water molecules play a major role in determining the

properties of the interphase at electrodes in contact with aqueous solutions, e.g. in terms of

the field-dependent orientation of H,O dipoles [ 1001 and their associated contributions to

the entropy of the interphase. In the presence of hydrated ions, which determine (cf.

Grahame, ref. 3) the dimensions of the compact part of the double-layer, interactions and

overlap with the hydration cosphere of the ion and the hydration coplane of the electrode

are important determinants of the microscopic structure, geometry and properties of thecompact double-layer region. In the case of ordered surfaces of single-crystals, it is likely

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that ordered, possibly commensurate structures of the coplane of water arise as has been

suggested by Wagner [217] and by Hubbard.

According to the infra-red work of Falk [218] for ClO,-, the H,O molecules in the

primary hydration shell are H-bonded to the 0-ligands of the Cl. There is some evidenceof the same effect in the case of SOd2- 219]. In the case of HClO, solution, the following

are possible sequences of the reactions on the surface: the tetrahedral ClO,-.n Hz0 anion

from the solution may lose x H,O upon adsorption giving ClO,-*(n-x) H,O adsorbed. For

H-bonded H,O, n must be at least equal to 4 as it is usually assumed that there is at least

one H-bond to each of the 0 ligands of the anion. However, enhanced binding of H,O to

pre-adsorbed 0, giving 8-10 H,O/O at low coverage of 0, has been observed upon

adsorption of H,O from the gas phase on a number of metals [220] and was considered a

general feature of such systems. These hydration numbers decrease with increase of the

coverage by 0. In the present case, the distance between the centers of 0 atoms in the

tetrahedral oxy-anions is such that one A-, without hydration H,O, could physically block

3 atoms of Au on the (111) plane and each hydration H,O will shield another 1 site. The

H,O associated with the top of the tetrahedral anion obviously will not influence the

number of sites required. The only change of site requirement will occur when, during

increase of 8, or during the charge transfer, the adsorbed anion will lose partially or

completely the H-bonded H,O molecules, a process that may cause rearrangement of the

overlay lattice of anions on the surface. Indeed, in the anodic i vs E profiles over the

double-layer potential region at Au, in HClO, solution a small double peak, which may be

attributed to rearrangement, arises when l/2 to l/3 of the anodic charge for this region has

been passed.

It will be interesting to estimate the number of sites shielded by one specifically

adsorbed ClO,-. In the light of results of Sass [221,222], calculation of this number can

be based on the supposition that the OH’s deposited in the reversible processes associated

with OAl and OA2, originate from the water molecules H-bonded to the 0 atoms of the

chemisorbed anions, rather than from patches of the free water network on the surface.

Then, the number of deposited OH’s, nOH,will be equal to the number of adsorbed

anions, nA, times the number of 0 atoms in an adsorbed anion in contact with the surface

(in the case of the ClO,- anions, 3n,) times the number of H,O’s H-bonded to each 0

atom, nhlo, i.e. n nH=3nAn,,,“. In terms of OH coverage, 8,,=8, n,,,O dn,“ where nEoand nE

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434 B.E. Conway

are the number of energy stages for deposition of OH geometrically expected and

experimentally observed, respectively. The excess charge passed over the double-layer

potential region is, as was described earlier, only q,.,=6-10 pC cm-*, which for 6=1

corresponds to 8,=0.08-0.14. The exact figure is difficult to evaluate because of the large

potential range (AE=0.9 V) of the adsorption and the uncertainty of the electrostatic

double-layer capacitance of this region. In the calculations, the value of 40 p.F cm-’ was

chosen. Owing to some uncertainty in the 9~ in the calculations of the hydration numbers,

an intermediate value of q*=lO PC cm-* or 0,=0.14 can be taken as the best choice.

The coverage of OH(‘-Y)m, eposited in the two energy states corresponding to the peaks

for the reversible processes OAl and OA2 is 0=0.3 (see Table IIl.8). A third energy state

for OH deposition from hydration-H,0 could be expected in HClO, solutions judging from

the symmetry of the (111) plane. The fact that it is not observed may be attributed to the

relatively weak chemisorption of ClO,- anion, causing the RTO process (OA3 peak) to

arise in HClO, solutions at potentials at which the peak for the third reversible stage

would otherwise have developed, i.e. the anions are desorbed before hydration H,O can be

oxidized to produce OH in the third energy state. Then for the total coverage for OH”-YF,

the coverage associated with this peak has also to be added: 8,,=0.34+0.17==0.15.

In the last case, n,/n,“=2/3 and from the above eqn. for t&n, n,,=3.6, i.e. the OHoY’-

originates from 3.6 H,O molecules H-bonded to each of the 3 0 atoms on the surface in

ClO,- giving 11 H,O’s per anion. This will shield 3+11=14 Au atoms per one complex of

anion with H-banded water. When the assumption is made that only 2 energy states exist,

as only 2 reversible stages are observed, nE/nEo=l and n,,=2.4. Then only 3+7=10 Au

atoms will be shielded.

The possibility of such a large hydration number seems to be supported by ESDIAD

results from gas phase adsorption of H,O on metals with pre-adsorbed 0 atoms [2 19-2221.

The oxygen-water interactions there may be viewed as B model system for the solvation of

ions on the surface [222]. At low coverages larger numbers of H,O’s can be H-bonded

(110).

III.9 Structural Changes at Au During Oxidation as Revealed by SERS

Spectroscopic techniques, when applicable, provide a very useful complement to

information generated from purely electrochemical experiments. Desilvestro and Weaver

[235] made SERS measurements (Figs. 49a and b) on the potential-dependent surface-

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oxygen vibrational bands observable at Au in aq. KOH and in HCIO, solutions up to OSV

or 1.3V vs RHE.

435

Fig. 49 a) SER spectra for gold in 1M KOH as a function of electrode potential.The spectra were recorded sequentially from bottom to top, using a

spectrometer scan rate of 1 cm.’ s-j. b) spectra of gold in 0.1 M NaClO, +0.01 M HClO,, as in (a), as a function of electrode potential.

III. 10 Mechanistic Overview of the Elementary Stages of Oxide Film Formation on

AU

The results obtained for the (111) plane, reviewed and discussed above, allow some

interesting improvements to be made to the scheme of the early stages of oxide film

growth and anion involvement, treated in Sections 111.1and 111.2,and generally (Part I) for

polycrystalline Au. They also provide the opportunity for giving a useful summary which

provides an overall picture of the early stages of oxide formation on the Au (111) single-

crystal plane. As will have been seen from the foregoing material, the process of surface

oxidation of noble metals is initially a complex multi-stage surface process, coupled with

parallel processes nvolving anion chemisorption and place-exchange (RTO), leading to

anion desorption and ultimately thicker film growth. In these respects, the mechanisms we

have developed, based on detailed analyses of results at single and polycrystalline Au,

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436 B.E. Conway

supersede, in important ways, earlier models based simply on the idea that successive

formation of supposed “MOH” and “MO” species arises with increasing anodic potential,

without reference to the “surface lattice” processes that must really be microscopically

involved.

At potentials somewhat positive to the P.Z.C. a reversible electrostatic adsorption of

hydrated anions first arises as on other metals:

xM+A y H,O+p H,O+M,,,A- z H,O+(x-3-z) M

where z=y+p. The anion adsorption takes place in a network of H,O on the surface by

abstracting H,O molecules and forming, with H,O, strongly H-bonded complexes similar

to what was suggested by Sass et al.[221,222] on the basis of their gas-phase

investigations mentioned earlier. The pseudocapacitance observed in the double-layer

potential region (Fig. 46) at more positive potentials (the D peaks) indicates that on the

Au (I 11) plane the anions are chemisorbed, i.e. they undergo partial or full charge-

transfer:

M3+,A-z H,O * M3+zA”-d)-* H,O+&-

It seems that 6=1 for HClO, solutions and 1.1 for H,SO, solutions.

Strong chemisorption is facilitated on the (111) plane by the compatibility of the three-

fold symmetry of Au atoms with that, axially, of the tetrahedral anions. The stronglychemisorbed anions cover the surface forming, due to short-range repulsive and long-range

attractive interactions through the metal, an overlay lattice or lattices of anions H-bonded

to H,O molecules. With increase of the coverage by the discharged anions, the hydration

number z will diminish:

M l+ZI-‘)-. zH,O+M,+,A”-“-v H,O+(z-v) M+(z-v) H,O

which may lead to a reconstruction of the overlay lattice giving rise to the sharp Dl peaks

(2~). Further discharge of anions will occur in the D2 peaks according to eqn. (10) butwith a smaller number, v, of hydration H,O’s, instead of z:

M,+,A-v H,O+M,+,A(‘-G’-~v H,O+&-

At more positive potentials, the partial discharge of the H,O H-bonded to the 0 atoms in

the anions starts to be possible and occurs on Au (111) in two energy states, usually in a

series of fast reactions (OAI, OA2 peaks):

M3+rA(‘-6)-* H,O+M,+,A”~s’~~(v- 1) H,OOH(“Y)s+H++ye--+

M3+,A(1-S)-.v-2) H,O*2 OH”-“‘-+H++ye-+...+

(50)

(51)

(52)

(53)

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M

3+2~O-sk. v OH(I -Y)-+H++~-

The experimental results that have been reviewed above suggest that the extent of charge

transfer for oxidation of the H,O on the (1 11) plane is -0.5.

v was found to be 7 for HClO, and, depending on concentration, 2-3 for H,SO,solutions. In this way, the OH seems to be incorporated into the overlay lattices of the

anions on the surface initially before those anions are eventually desorbed. The results

suggest that the effect of the specifically adsorbed anions is not only to hinder the

oxidation of Au by blocking the surface for discharge of H,O from the free water network,

but the anions also facilitate partial discharge of H,O H-bonded to 0-ligands of the Cl or

S in the anions providing, in this way, a new reaction path for the discharge of H,O which

occurs at less positive potentials than the main oxidation reaction in the RTO process.The potential range in which process [47] takes place is such that the surface freed by

anions in that reaction can immediately be further oxidized:

M+H,O+MOH+H’+e-,

increasing the cutrents in the RTO peaks.

The MOH dipoles may form on an anion-free surface 2 or 3 sublattices. These

correspond to the most stable states of MOH as the latter is reduced in the most negative

peak or peaks (the OC2’s). However, the major fraction of the MOH dipoles undergoesreconstruction in the turn-over process leading to phase formation:

HHH H H000 -3 OMOMMM MOM

H

This material. which is reduced in the OC3 peak, prevails when 8,” incretlses.

The passage of the 2nd electron, MOH -+ MO+H’ + e-, which produces a

pseudo-2-dimensional phase of orderly “AuO”, arises in the OA4 potential region, giving a

clearly distinguished peak with an “half-width” of only 140 mV. As the desorption of

anions is gradual on the Au (111) plane, and therefore the RTO process occurs

simultaneously with the first and partially the second electron transfer, the (111) plane is

evidently not the most suitable for investigations of the RTO process itself.

The processes that are presented above can be conveniently shown in summary form as

was shown in Scheme 2.

(54)

(55)

(56)

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438 B.E. Conway

It should be remarked, in conclusion, that for Au, there is, of course, no opportunity of

indirectly following surface structural changes due to reconstruction by an in situ UPD H

determination, as is possible at Pt. However, LEED or RHEED surface structure analyses

[227] do indicate changes of surface structure at Au( 111) and (100) due to reconstruction,

as is also directly observable under STM or AFM (see following section).

III. 11 STM Observations of Topographical Changes at Oxidized Au Surfaces

Some difficulties are encountered in obtaining atomic-resolution STM images of noble-

metal surfaces bearing oxide films or adsorbed anions. However, Ullmann, Will and Kolb

[244] have been able to make significant observations on surface topography of Au( 111)

in relation to oxidation following desorption of UPD Cu. Several of their results are

shown below as STM image pictures.

An STM image (100 x 100 nm2) of Au( I1 1) in sulfuric acid solution at a potential where

the surface is bare of Cu. Monoatomic high steps, crossing each other at an angle of 60”,

and atomically flat terraces are seen. The 6 x 6 nm2 zoom shows the individual gold

atoms on the surface.

When the potential is stepped into the oxide region (here: +1.3 V vs SCE), the image is

somewhat blurred due to the oxide. Importantly, reducing the oxide by a potential step

back to +0.5 V shows that monoatomic high (deep) Au islands and holes are generated.

Waiting at that potential for a few minutes causes the islands to disappear (electrochemical

annealing) and even the less mobile holes grow (Figs. 50. 51 and 52).

Scanning the potential slowly up to 1.3 V and back also roughens the surface by creating

monoatomic deep holes, but their counterparts, the islands, are not present. Holes and

islands appear simultaneously when scanning the potential fast. This means that during

slow reduction the adatoms have enough time to diffuse on the surface and be

incorporated at steps: hence, no islands arise. It is known from other studies, that the

surface mobility of gold atoms is greatly enhanced at the positive end of the double-layer

region due to adsorption of anions at steps. If you are long enough in that region on the

way to +OS V, because of slow scan, gold islands disappear due to electrochemical

annealing. After 2-3 potential cycles, channels are formed between neighboring holes,

indicating a non-unidirectional surface corrosion. At present it is not certain whether the

STM tip is interfering or not.

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500 mV vs. Cu/Cu’+

before

200 run x 200 nm

Fig. 50 STM images for Au( 111) in 0.05 M H,SO, + 0.1 mM CuSO, showingchanges of topography caused by oxidation at 1.30 V vs CuKu”; potentialstepped from 0.50 to 1.30 back to 0.50 V at indicated times. (FromUllmann et al., ref. 244).

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440 B.E. Conway

Au( 111) / 0.05 M l&SO, + 0.1 n&I CuSO,

Atomic distance:

d u*a = 2.9 ii

6 nm x 6 nm

E = +500 mV vs. Cw’Cu”

E

0”Isi 100

wri

0-----------25

25 --r-----50 _.__. -T----.-.

75 I 1___ 100”

ret

Fig. 51 STM images of Au(l11) in 0.05 M H,SO, (with 0.1 mM C&O, at 0.50 V

vs Cu/Cu” potential. Interatomic distance Au-Au = 0.29 nm. (From

Ullmann et al., ref. 244).

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holes

Fig. 52 STM image of Au( 100) in O.lM HClO,; left, at 0.8 V after cycling

between 0 and 0.8 V En: right, at 0.8 V after one cycle between 0 and 1.5V En. Note generation of island-hole structure. (From Itaya and Honbo,ref. 231).

These results show principally that roughening of Au surfaces depends much more

decisively on the history of oxidation-reduction cycling than on the holding time at the

oxidation potential. This is consistent with various observations in the literature, e.g. ref.

26, that potential cycling at Pt and Au is most effective in generating thick oxide films

accompanied by extensive surface roughening, as is also confirmed by LEED.

In other recent STM work on Au(100) by Kolb et al. [244], it has been found that even

over the double-layer potential region, topographical features are generated with change of

potential in one direction and are removed upon reversal of the direction of potential

change (Fig. 53). These effects evidently are very sensitive to changes of surface-excess

Gibbs energy. One is reminded of the substantial variation of that quantity at the Hg

electrode where variations of ?12% around the P.Z.C. for changes of ea. +OSV can arise.

At a solid electrode of relatively low lattice energy, related effects are evidently

manifested with a consequent modification of interfacial toppology, prior to onset of

surface oxide formation, though chemisorbed anions may be involved.

Similar indications of changes of surface topology at Au caused by changes of potential

in the “double-layer” region have been observed in recent AFM work from CNRS,

Meudon, by Clavilier and Hamelin.

In an important recent paper [249], Weaver et al. have studied the adsorption of sulfate

ion on Au( 111) by means of IR spectroscopy and STM. An adlayer of sulfate ions is

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442 B.E. Conway

formed at a fractional surface coverage of ca. 0.2, confirming the conclusions reached in

refs. 56 and 57 from analysis of CV curves, and having a (J3 x J7) structure. These

results are also consistent with those derived from coulometric and radiotracer

(Wiechowski) data. The assignment that can be made to the observed “S-O” stretching

band at 1155-1230 cm-‘, baed on its insensitivity to electrolyte conditions, e.g. pH, leads to

the conclusion that it is the sulfate ion rather than the bisulfate ion that is preferentially

adsorbed despite the measurements having been carried out at pH-1 which, in relation to

the pK,,, of H,SO, in water would suggest an appreciable concentration of HSO;

adsorption seems to be preferred according to recent thermodynamic analysis of double-

layer charge-potential relations by Lipkowski et al. [250].

Related to this matter is the solvation factor in ionic chemisorption [251,252], on this

basis of which ions having weaker hydration (HSO, < SO,%) are expected to be more

chemisorbable (through lone-pair donicity) than corresponding ones having larger solvation

energies.

Fig. 53 STM images for Au( 100) in 1M H,SO, showing changes of surfacetopography resulting from potential variation over the double-layer chargingregion. (From ref. 244).

A model of the SO,,” anion + water adsorption at Au( 11 ), based on the STM results, is

shown in Fig. 54 from the paper of Weaver et al. [249]. An height-shaded STM image of

the Au( 1 11) surface at 0.9V, SCE is shown in Fig. 55 [249]. At higher potentials

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(E> 1 OV) the SO,’ coverage decreases but the local (J3 x J7) STM pattern remains,

microscopically, but due to extension of surface oxidation of the Au, the oxide film

progressively replaces (cf. refs. 56 and 57) the regions of chemisorbed SO,*-, i.e. domain

rather than partially open array surface structures are involved.

443

Au(lll)-SO4 (43x1/7

Fig. 54 Suggested real-space (J3 x J7) structure of SO,,‘- adlayer on Au( 111).Larger hatched circles represent SO,‘- ions, smaller ones possiblecoadsorbed cations and/or oriented hydration water. (From Weaver et al.,ref. 249)

Fig. 55 30” off normal height-shaded STM image for Au( 111) in 0.1 mol dme3aq.H,SO, at 0.9 V, SCE. (From Weaver et al., ref. 249).

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444 B.E. Conway

Part IV: SOME ASPECTS OF THE SURFACE OXIDATION OF Rh

The surface oxidation behavior of the element Rh is in many ways more similar to that

of Pt than of Au or Pd. Relatively few works have appeared in recent years on the

surface electrochemistry of this metal but Jerkiewicz has carried out some significant

studies recently.

The question of applicability of the Mott-Cabrera theory [ 1171 (inverse log growth

kinetics) to noble-metal oxide film growth has been further examined by Jerkiewicz and

Borodzinski [245,246] for the case of polycrystalline Rh. It appears that linear relations

between qoX-’and the growth time, t,, arise but only for appreciable extents of film growth.

For small extentsof film formation, the inverse log law does not apply and, in fact, the

direct law as found at Pt [ 142,157] may better represent the results. However, significant

differences in the film growth behavior at Rh in comparison with Pt were noted. At Rh, a

transition from “direct” to “inverse” log growth kinetics is indicated when a nominal

Rh(OH), coverage of I 1 ML has been reached and exceeded. The electron transfer

through the oxide film from H,O to the metal substrate proceeds by tunneling (cf. refs.

248 and 5X) and is the rate-controlling step in the oxide film growth [247] after its initial

stages have been completed.

Acknowledgments

Grateful acknowledgment is made to the Natural Sciences and Engineering Research

Council of Canada for support of this work and of earlier research on which parts of this

review are based. Special thanks are due to Drs. H. Angerstein Kozlowska (previously of

this laboratory), A. Hamelin (CNRS, Meudon) and J. Clavilier (NRS, Meudon) for many

discussions on the topic of this article and for their contributions to jointly authored

papers. Similar acknowledgment is due to Dr. B.V. Tilak (Occidental Chemical

Corporation, Buffalo, N.Y.). Grateful acknowledgements are also made to various authors

of original papers in the field for permission to reproduce diagrams and, in some cases, for

useful discussions of their results and provision of recent results, e.g. from Professor

D.M. Kolb and Professor P.N. Ross.

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R. Woods in Electroanalytical Chemistry, Vol. 9, Ed. A.J. Bard, pp. 2-144, MarcelDekker, New York (1976).M.J. Dignam, Chapter 5 in Comprehensive Treaty of Electrochemistry, Vol 4, Eds.J.O’M. Bockris, E. Yeager and B.E. Conway, Plenum, New York (1986).B.E. Conway, Prog. in Surface Science, 16, 1 (1984).L.D. Burke, Chapter 4 in Modem Aspects of Electrochemistry, Vol 18, Eds.J.O’M. Bockris, R.E. White and B.E. Conway, Plenum Publ. Corp. New York

(1986).B.E. Conway and B.V. Tilak, Adv. Catalysis, 38, 1 (1992).G.P. Sakelleropoulos, Adv. Catalysis, 39, 217 (1981).H.A. Laitinen and C.G. Enke, J. Electrochem. Sot., 107, 773 (1960).F.G. Will and C.A. Knorr, Zeit Elektrochem., 64, 258 (1960).B.E. Conway, H.A. Kozlowska, W. Sharp and E. Criddle, Anal. Chem., 45, 1331(1973).S.W. Feldberg, C.G. Enke and C.E. Bricker, J. Electrochem. Sot., 110, 826 (1963).

F.C. Anson and J.J. Lingane, J. Amer. Chem. Sot., 79, 4901 (1957).W.G. French and T. Kuwana, J. Phys. Chem., 68, 1279 (1964).A. Hickling, Trans. Faraday Sot., 41, 333 91945).

J.S. Mayell and S.H. Langer, J. Electrochem. Sot., 111, 438 (1964); ibid 113, 385(1966).

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S. Schuldiner and T.B. Warner, J. Electrochem. Sot., 112, 212 (1965) and 113, 573(1966).

T.V. Kalish and R. Burshtein, Dokl. Acad Nauk SSSR, 81, 1093 (1951); 88, 863

(1953).

J.A. Davies, D.P. Jackson, J.B. Mitchell, P.R. Norton and R.L. Tapping, Nut. Inst.Meth., 132, 609 (1976).K.J. Vetter and D. Bemdt, Zeit. Elektrochem., 62, 378 (1958).J.P. Hoare, J. Electrochem. Sot., 109, 858 (1962); 110, 245 (1963); 111, 232

(1964); 111, 610 (1964); 111, 988 (1964); 112, 602 (1965); 112, 608 (1965); 112,

849 (1965); 112, 1129 (1965).J.P. Hoare, Electrochim. Acta, 11, 311 (1966); 11, 549, (1966); 12, 260 (1966).

J.P. Hoare, J. Electrochem. Sot., 110, 1019 (1963); see also 77zeElectrochemistryof Oxygen, Wiley-Interscience, New York (1968).J.O’M. Bockris and A.K.M.S. Huq, Proc. Roy. Sot., London, A237, 277 (1956).M. Rao, A. Damjanovic and J.O’M. Bock& J. Phys. Chem., 67, 2508 (1963).J.O’M. Bockris, A.K.N. Reddy and M.A. Genshaw, J. Chem. Phys., 48, 67 1(1968).

B.E. Conway and D.M. Novak, J. Electroanal. Chem., 99, 133 (1979).B.E. Conway, B. MacDougall and H. Angerstein-Kozlowska, Proc. 2nd Intl.Conference on Fuel Cells, Bruxelles (1968), S.E.R.A.I., Bruxelles, (1969).A.K. Vijh and B.E. Conway, Chem. Rev., 67, 623 (1967).

H. Angerstein-Kozlowska, B.E. Conway and W.B.A. Sharp, J. Electroanal. Chem.,43, 9 (1973).

H. Angerstein-Kozlowska, B.E. Conway and A, Hamelin, Electrochim. Acta, 31,1051 (1986).

H. Angerstein-Kozlowska, B.E. Conway, A. Hamelin and L. Stoicovicu, J.Electroanal. Chem., 228, 429 (1987).K.J. Vetter and J.W. Schultze, J. Electroanal. Chem., 34, 131 (1972).K.J. Vetter and J.W. Schultze, J. Electroanal. Chem., 34, 141 (1972).W. Lorenz and G. Sal%, Zeit. Phys. Chem. Leipzig, 218, 259 (1962).J.W. Schultze and F.D. Koppitz, Electrochim. Acta, 21, 327; 337 91976).R. Balaschova and V.E. Kazarinov, J. Electroanal. Chem., 3, 135 (1969).

S. Bamartt, J. Electrochem. Sot., 106, 722 (1959).J.J. MacDonald and B.E. Conway, Proc. Roy. Sot., London, A269, 419 (1962).M. Pourbaix, Atlas of Electrochemical Equilibria, Cebelcor, Bruxelles, Gauthier-Villars, Paris ( 1963).J.A. Davies et al., see ref. 44.L. Young, Anodic Oxide Films, Academic Press, New York (1961).S.E.S. El Wakkad and S.H. Emara, J. Chem. Sot., London, 461 (1952); 3094(1954).

S. Bamartt, J. Electrochem. Sot., 106, 991 (1959).R. Parsons and W. Visscher, J. Electroanal. Chem., 36, 329 (1972); see also M.Barrett and R. Parsons, Disc. Faraday Sot., p.72 (1970).For example, see B.E. Conway, Theory and Principles of Electrode Processes,

37.

38.

39.

40.

41.

42.

43.

44.

45.

46.

47.

48.

49.

50.

51.

52.

53.

54.

55.

56.

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59.

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71.

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74.

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78.

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83.

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89.

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92.93.

94.95.

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103.

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M. Fleischmann, P.J. Hendra and A.J. McQuillan, Chem. Phys. Letters, 26, 163(1974).A. SevCik, Coll. Czeck. Chem. Comm., 13, 349 (1948).S. Srinivasan and E. Gileadi, Electrochim. Acta, 11, 321 (1966).H. Angerstein-Kozlowska, B.E. Conway and J. Klinger, J. Electroanal. Chem., 75,45 (1977); 75, 61 (1977).P. Stonehart, H. Angerstein-Kozlowska and B.E. Conway, Proc. Roy. Sot., London,A310, 541 91969).H. Angerstein-Kozlowska and B.E. Conway, J. Electroanal. Chem., 95, 1 (1979).F.P.Bowden and E.K. Rideal Proc. Roy. Sot., London, Al25, 446 (1929).J.A.V. Butler, Electrocapillarity, Chapter 8, Methuen, London (1940).B.V. Erschler, M. Deborin and A.N. Frumkin, Acta Physiocochim. U.R.S.S., 8, 565(1938).J.F. Armstrong, S. Himsworth and J.A.V. Butler, Proc. Roy. Sot., London, A137,604 (1932); A143, 89 (1933).H. Angerstein-Kozlowska and B.E. Conway, J. Electroanal. Chem., 7, 109 (1964).A.L. Hodgkin and A.F. Huxley, Nature 144, 710 (1939); see also J. Physiol., 116,424 (1952).A. Hickling, Trans. Faraday Sot., 38. 27 (1942).B.E. Conway and E. Gileadi, Trans. Faraday Sot., 58, 2493 (1962).B.E. Conway, B. MacDougaU and H. Angerstein-Kozlowska, Disc. Faraday Sot.,

56, 199 (1973).A.T. Hubbard, R.M. Ishikawa and J. Katekaru, J. Electroanal. Chem., 86, 27 1( 1978).K. Engelsman, H. Schmidt and J.W. Lorenz, J. Electroanal. Chem., 114, 1 (1980).A.A. Michri, A.G. Pshchenichnikov and R.K. Burshtein, Elektrokhimiya, 8, 364(1972).T. B iegler and R. Woods, J. Electroanal. Chem., 20, 73 (1969); 21, 457 (1969).See, e.g. G. Somorjai, Principles of Surface Chemistry, Prentice Hall, EnglewoodCliffs, N.J. (1972).J. Clavilier, J.Electroanal. Chem., 107, 205 (1980); 100, 401 (1978).B.E. Conway, N. Marincic and J. Wojtowicz, J. Chem. Phys., 48, 4333 (1968).

M.W. Breiter, Electrochim. Acta, 7, 533 (1962); 8, 925 (1963).V.S. Bagotzky, Y.B. Vasil’yev, J. Weber, J.N. Pirtskhaleva and B. Knots, J.Electroanal. Chem., 27, 31 (1970).B.E. Conway and D.M. Novak, J. Chem. Sot., Faraday Trans. I, 77, 2341 (198 1).V. Gutmann, Chemistry in Britain, 7, 102 (1971).J. O’M. Bockris, M.A.V. Devanathan and K. Mtiller, Proc. Roy Sot., London,A274, 55 (1963); see also R.J. Watts-Tobin and N.F. Mott, Phil Mag., 6, 133(1961).J.O’M. Bockris and T.N. Anderson, Electrochim. Acta, 9, 347 91964).A.A. Kornyshev, AI. Rubinshtein and M.A. Vorotyntsev, Phys. Stat. Solidi, (b)84,

125 (1977).P.H. Cutler and J.C. Davis, Surf. Sci., 1, 194 (1964).

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104.105.

106.

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108.109.110.

111.112.113.114.115.

116.117.118.119.120.121.122.123.

124.

125.

126.

127.

128.

129.

130.

131.132.133.134.

135.136.

137.

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W. Visscher and J.M. Otten, J. Electroanal. Chem., 55, 1; 13 (1974).J.L. Ord and F.C. Ho, J. Electrochem. Sot., 118, 46 (1971).J.L. Ord and D.J. de Smet, J. Electrochem. Sot., 116, 762 (1969).

M.A. Barrett and R. Parsons, Symposium of the Faraday Sot., 4, 72 (1971).B.E. Conway, Symposium of the Faraday Sot., 4, 86 (1971).T. Biegler and R. Woods, J. Electroanal. Chem., 20, 73 (1969).G. Jerkiewicz and B.E. Conway, J. Electroanal. Chem., 334, 359 (1992).G. Sauerbrey, Z&t. Phys., 155, 206 (1959).

R. Schumacher, G. Borges and K. Kanazawa, Surf. Sci., 163, 621 (1985).S. Bruckenstein and M. Shay, J. Electroanal. Chem., 188, 131 (1985).S. Bruckenstein and M. Shay, Electrochim. Acta, 30, 1295 (1985).

W. Stiickel and R. Schumacker, Ber. Bunsenges. Phys. Chem., 91, 345 (1987).S. Bruckensetin and S. Swartirajan, Electrochim. Acta, 30, 851 (1985).K. Kanazawa, 0. Melroy, D. Bultry and J. Gordon, J. Amer. Chem. Sot., (inpress); Anal Chem., 57, 1770 (1985).S. Gilman in Electroanalytical Chemistry, Ed. A.J. Bard, Vol. 2, Marcel Dekker,New York (1967).M. Barrett and R. Parsons, J. Electroanal. Chem., 42, App. 1 (1973).M.W. Breiter, J. Electroanal. Chem., 7, 38 (1964).S. Gilman, Electrochim. Acta, 9, 1025 (1964).L. Formaro and S. Trasatti, Electrochim. Acta, 12, 1457 (1967).

A.J. Arvia, Acta Cient. Venez., 31 496 (1980).A. Wieckowski, P. Zelenay and K. Varga, J. Chim. Phys., Chim. Biol., 88, 1247(1991).S. Schuldiner and T.B. Warner, J. Phys. Schem., 68, 1223 (1964)(.A. Horanyi, Bull. Electrochem., 5, 235 (1989); see also G. Horanyi and E.M.Rizmayer, J. Electroanal. Chem., 218, 337 (1987).H. Angerstein-Kozlowska, K. Tellefsen and B.E. Conway, J. Electroanal. Chem.,24, 1045 (1989).

140. D.H. Everett, Trans. Faraday Sot., 50, 187 (1954); 51, 1551 91955).141. D.H. Everett and P. Nordon, Proc. Roy. Sot., London, A259, 351 (1960).142. D. Gilroy and B.E. Conway, Can. J. Chem., 46, 875 (1968).

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R. Simpraga and B.E. Conway, J. Electroanal. Chem., 355, 79 (1993).M.J. Dignam and D.B. Gibbs, Can. J. Chem., 48, 1242 (1970).B.E. Conway and J. Mozota, Electrochim. Acta., 28, 1 (1983).G. Tremiliosi-Filho, G. Jerkiewicz and B.E. Conway, J. Electroanal. Chem., 297,435 (1991).

S. Hadzi-Jordanov, H. Angerstein-Kozlowska, M. Vukovic and B.E. Conway, J.Electrochem. Sot., 125, 1471 (1978).S. Shibata, Bull. Chem. Sot., Japan, 36, 53 (1963); 37, 12 (1964); 40, 696( 1967). See also ref. 195.S. Shibata and M. Sumino, Electrochim Acta, 20, 739 (1975).W. Visscher and M.A.V. Devanathan, J. Electroanal. Chem., 8, 127 (1964).D. Dickertmann, J.W. Schultze and K.J. Vetter, J. Electroanal. Chem., 55, 429(1974).A. Wieckowski, in Modern Aspects of Electrochemistry, 21, 54 (1990); see alsoref. 136.A.N. Frumkin, Sov. Phys., 4, 239 91933).

L. Pauling, The Nature of the Chemical Bond, Cornell Univ. Press (1930).B.E. Conway, B.V. Tilak, B. Barnett and H. Angerstein-Kozlowska, J. Chem.Phys., 93, 8361 (1990).B.E. Conway and G. Jerkiewicz, J. Electroanal. Chem., 334, 359 (1992). See alsoM. Farebrother, M. Goledzinowski, ibid, 291, 469 (1991).G. Tremiliosi-Filho, G. Jerkiewicz and B.E. Conway, Langmuir, 8, 658 (1992).H. Dietz and H. Giihr, Electrochim. Acta, 8, 343 (1963).J.W. Schultze and K.J. Vetter, Ber. Bunsenges, Phys. Chem., 75, 470 (1971).V.S. Bagotzki, W.J. Lukjanytschewa, A.J. Osche and W.J. Tichomirov, Dokl.Akad. Nauk, SSSR, 159, 444 (1964).

M.W. Breiter, Electrochim. Acta, 7, 533 (1962); 8, 925 (1963).L. Formaro and S. Trasatti, Electrochim. Acta, 12, 1457 (1967).M.W. Breiter, Electrochim. Acta., 7, 25 (1962).S.B. Brummer and AC. Makrides, J. Phys. Chem., 68, 1448 (1964).B.E. Conway, N. Marincic, D. Gilroy and E.J. Rudd, J. Electrochem. Sot., 113,

1144 (1963).V.I. Tikhomirova, A.I. Oshe and V.S. Bagotzky, Dokl. Akad, Nauk, S.S.S.R., 159,644 ( 1964).T. Biegler, D.A. J. Rands and R. Woods, J. Electroanal. Chem., 35, 209 (1972).J. Clavilier, J. Electroanal. Chem., 107, 211 91980); 107, 205 (1980); 127, 281(1981); 263, 109 (1989).

A.J. Appleby, J. Electrochem. Sot., 120, 1205 (1973).B.V. Tilak, Angerstein-Kozlowska and B.E. Conway, J. Electroanal. Chem., 48, 1(1973).T. Biegler, D.A.J. Rand and R. Woods, J. Electroanal. Chem., 29, 269 (197 1); 47,353. 1973).S. Shibata, J. Electroanal. Chem., 89, 37 (1978).W. Visscher and M. Blijlevens, J. Electroanal. Chem., 47, 363 (1973); see alsoJ.M. Otten and W. Visscher, J. Electroanal. Chem., 55 1;13 (1974).Yu.Ya Vinnikov, V.A. Shepelin and V.I. Veselovskii, Electrokhim, 9, 552; 649;1557 (1973).

Yu.Ya. Vinnikov, V.A. Shepelin and V.I. Veselovskii, Elektrokhim., 8, 1229; 1384,(1972).

144.145.146.14x.

149.

150.

151.152.153.

154.

155.

156.

157.

158.

159.160.161.162.

163.164.165.166.167.

168.

169.170.

171.172.

173.

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176.

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R. Woods, Chapter 1 in Electroanalytical Chemistry, Vol. 9, Ed. A.J. Bard, MarcelDekker, New York (1976).S.D.James, J. Electrochem. Sot., 116, 1681 (1969).L.D. Burke and M.B.C. Roche, J. Electroanal. Chem., 137, 175 (1982); see alsoL.D. Burke, Chapter 4 in Modern Aspects of Electrochemistry, Vol. 18, Eds. R.E.

White, B.E. Conway and J. O’M. Bockris , Plenum, New York (1987).J. Balej and 0. Spalek, Coll. Czech. Chem. Comm., 37, 499 (1972).M. Faraday, Pogg. Ann., 33, 438 (1834); see A.J. Vijh and B.E. Conway, Chem.Rev., 67, 623 (1967).H. Kolbe, Ann., 69, 257 (1849); ibid, 113, 125 (1860).

D.M. Novak and B.E. Conway , J. Chem. Sot. Faraday Trans. I, 77, 2341 (1981).V.S. Bagotzky and Y.B. Vasil’yev, J. Electroanal. Chem., 27, 31 (1970).K.S. Kim, N. Winograd and R.E. Davis, J. Amer. Chem. Sot., 93, 6296 (1971).D.A.J. Rand and R. Woods, J. Electroanal. Chem., 35, 209 (1972).Y.M. Tyurin and G.F. Volodin, Elektrokhimiya, 6, 1186 (1970).V.E. Kazarinov, Elektrokhimiya, 2, 776 (1966).

B.E. Conway, G. Tremiliosi-Filho, and G. Jerkiewicz, J. Electroanal. Chem., 297,1435 (1991).D.A.J. Rand and R. Woods, J. Electroanal. Chem., 35, 209 (1972).J.S. Hammond and N. Winograd, J. Electroanal. Chem., 78, 55 (1977); see also J.Amer. Chem. Sot., 93, 6296 (1971).M. Peuckert, Electrochim. Acta, 29, 1315 91984).G.C. Allen, P.M. Tucker, A. Capon and R. Parsons, J. Electroanal. Chem., 50, 335( 1974).S. Shibata, Electrochim. Acta, 22, 175 (1977).V.I. Birss and A. Damjanovic, J. Electrochem. Sot., 130, 1688 (1983); 130, 1694

(1983).M Farebrother, M. Goledzinowski, G. Thomas and V.I. Birss, J. Electroanal.Chem., 297, 469 (1991).S. Gottesfeld, M. Yaniv, D. Laser and S. Srinivasan, J. de Physique, Colloque C5,suppl., 11, 38, 145 (1977).S.G. Roscoe and B.E. Conway, J. Electroanal. Chem., 224, 163 (1987).S. Shibata and M.P. Sumino, Electrochim. Acta., 17, 2215 91972).W.B.A . Sharp, Ph.D. Thesis, University of Ottawa (1974).S. Gottesfeld, J. Electrochem. Sot., 127, 1922 (1980); 126, 742 (1979).S. Gottesfeld in Electroanal. Chem., Vol. 15, Ed. A.J. Bard, Marcel Dekker, NewYork (1989).

S. Gottesfeld and S. Srinivasan, J. Electroanal. Chem., 86, 89 (1978).S. Gottesfeld, G. Maia, J.B. Floriano, G. Tremiliosi-Filho, E.A. Ticianelli and E.R.Gonzalez, J. Electrochem. Sot., 138, 3219 (1991).B.D. Caban, J. Horkans and E. Yeager, Surf. Sci., 37, 559 (1973).T.A. Delchar and F.C. Tompkins, Proc. Roy. Sot., London, A300, 141 (1967).A. Hamelin and M. Sottto, C.R. Acad. Sci. (Paris), 271C, 609 (1970).M. Sotto, J. Electroanal. Chem., 69, 229 (1976); 70, 291 (1976); 72, 287 (1976).A. Hamelin, J. Electroanal. Chem., 138, 395 (1982).J. Clavilier and C. Nguyen van Huong, J. Electroanal. Chem., 80, 101 (1977).A. Hamelin, Modern Aspects of Electrochemistry, Eds. B.E. Conway,J. O’M. Bockris and R.E. White, Vol. 16, Chapter 1, Plenum, New York 91985).R.K. Burshtein, Elektrokhimiya, 3, 349 (1967).

450

178.

179.180.

181.182.

183.184.185.186.187.188.189.

190.

191.192.

193.194.

195.196.

197.

198.

199.200.201.202.203.

204.205.

206.207208.209.210.211.212.

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214.215.216.

217.218.219.

220.221.222.223.

224.225.

226.

227.

228.229.230.

231.

232.233.234.

235.236.237.23X.

239.240.

241.242.

243.244.

245.246.247.

248.249

P.N. Ross, 187th A.C.S. Meeting. St. Louis (1984).B.E. Conway and H. Angerstein-Kozlowska. J. Electroanal. Chem., 113, 63 (1980).C. Nhuyen van Huong, C. Hinnen and J. Lecoeur, J. Electroanal. Chem, 106. 185(1980): see also C. Nguyen van Juon, C. Hinnen, J. Lecoeur and R. Parsons, J.Electroanal. Chem., 92, 239 (1978).

F.T. Wagner and P.N. Ross, Appl Surf. Sci., 24, 87 (1985).G. Brink and M. Falk, Can. J. Chem., 48, 3019 (1970).B.E. Conway, Ionic Hydration in Chemistry and Biophysics, Elsevier, New York,(1981).T.E. Madey and F.P.Netzer, Surf. Sci., 17, 549 (1982).K. Bange, D. Grider and J.K. Sass. Surf. Sci., 126, 437 (1983).K. Bange, D.E. Grider, T.E. Mady and J.K. Sass, Surf. Sci., 137, 38 (1984).H. Angerstein-Kozlowska, B.E. Conway, A. Hamelin and N. Stoichovicu, J.Electroanal. Chem., 228, 429 ( 1987).D. Aberdam, R. Durand, R. Faure and F. El-Omar, Surf. Sci., 171, 303 (1986).

P.N. Ross, J. Electrochem. Sot., 126, 67 ( 1979).

F.T. Wagner and P.N. Ross, J. Electroanal. Chem., 150, 141 (1983) and 250, 301(1988).P.N. Ross in Structure of Electrified Interfaces, J. Lipkowski and P.N. Ross, eds.,Frontiers of Electrochemistry Series, VCH, New York and Wienhelm, (1993).A. D’Agostino and P.N. Ross, Surf. Sci., 185, 88 (1987).

F. Wagner and P.N. Ross, Surf. Sci., 160. 305 9 (1985).K. Itaya, S. Sugawara, K. Sashikata and N. Furuya, J. Vat. Sci. Technol., AS, 341(1990); see also, ibid, B9, 457 (1991).T. Hachiya, H. Honbo and K. Itaya, J. Electroanal. Chem., 315, 275 (1991); seealso J. Chim. Phys., 88, 1477 (1991).

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