On the electrochemical behavior of H2O2 AT Ag in alkaline solution

ON THE ELECTROCHEMICAL IN ALKALINE

BEHAVIOR SOLUTION

OF H,02 AT Ag

MASATAKA HONDA Hokkaido University of Education

at Kushiro,

Kushiro

085, Japan

and

TAKURO KODERA Research Institute

for Catalysis,

Hokkaido

University,

Sapporo

060, Japan

and HIDEAKI Department

of Chemistry,

ICITA

Faculty of Science, Hokkaido

University,

Sapporo

060, Japan

(Received 22 June 1982; in revised form 20 October 1982) Abstract-Etectrochemical lo-‘4.3

oxidation

M H,O, and 2 x lo-‘-1.0

and reduction of H,O, on Ag were studied in alkaline solution of M KOH under N2 bubbling. Steady i-4 cwves obtained by a cyclic

potential sweep method in a potential range where no electrode oxidation takes place, lead to the following results: (1) is (Acm-*) (cathodic limiting current density) = 1.0 x [H202];” (M], (2) ia (Acm-‘) (anodic limiting one) = i:([KOH] B [H,O,]=) or 1.0 x [KOH]L~o (M) ([KOH] < [H,O,],). (3) 4, (V) (mixed potential) = O.L26&.060 10g[K0H]‘~~ and (4) (a&/&)+ _ .+m( Rcm2) (reaction resistance at @ = $,) = 0.057 x [H,O&‘,*

(M-l),

where [H2021T

designates

a total H,O* concentration

and the others have

their usual meanings. The above results are explained by the following mechanism, HO; formed by the reversible chemical reaction, H,02 -t OH- + HO; fH,O, is oxidised in anodic reaction by two steps: HO; AHO, +e- and HO,(a)+OH-+O,+H,O+e-, whereas in cathodic reaction, H,O, is reduced by H,O, +e- _PrzOH(a) +OH-, OH(a)+e-+ OH-. Here, + designates a rate determining step. Catalytic decomposition of H,Oz on the electrode is also discussed.

INTRODUCTION

EXPERIMENTAL

Several studies have been reported regarding to the electrochemical[14] and chemical[l, 2,5, 61 (catalytic decomposition) reactions of H,02 on Ag in alkaline solution. Regarding the latter reaction, there is a contradiction in reported mechanisms: one of the groups[2,5,7] has asserted the existence of a chemical mechanism and the other group[l] has insisted that the mechanism consists of only electrochemical steps. Another contradiction is observed as regards the potential of Agat open circuit. This potential is tailed a mixed potential from counteracting electrochemical reactions on one and the same electrode. Thus, the oxidation of H,O, and the reduction of surface oxide[3,7] or the oxidation and reduction of HIO,[l] has been proposed. This paper deals with the behavior of the open circuit potential and the cathodic and anodic polarisation curves of Ag in O2 free alkaline H,O, solutions and aims to clarify the mechanism of the electrochemical oxidation and reduction of H,O,. Catalytic decomposition of HzOl is also discussed in terms of the electrochemical steps.

Silver wire of 99.99’j, from Johnson Mathey Co. served as a test electrode which had an apparent surface area from 0.1 to 0.3 cm2 (dia. 0.10 cm). The electrode was polished with emery paper and then washed with benzene, methanol and conductivity water. Bright Pt foil was used as a counter electrode (area ca. 2 cm’) and placed together with the test electrode in a cell made of Pyrex glass. The reference electrode was a mercury-mercuric oxide electrode in blank solution without H,Oz. Solution contact was made through a closed moistened tap in order to avoid mixing of both the solutions. The potential of the reference electrode was checked against reversible hydrogen electrode before each measurement. Solutions were prepared from reagents of special grade and conductivity water. Concentration of KOH was 2 x 10m3-1 M. At low KOH concentrations, NaClO, was added to reduce a solution resistance. Concentration of H,Oz was varied from 10e3 to 0.3 M by adding successively small amount of stock H20, solution of known concentration at a given KOH 727

728

MASATAKA

HONDA,TAKUROKODERAAND

concentration under the bubbling of purified N, gas. Concentration of H,O, was determined by a permanganate titration method before and after the experiments. Measurements were carried out by a cyclic potential sweep method (potentiostat; Nichia Keiki Co. HP-500, function generator; Nichia Keiki Co. S-SA, and X-Y recorder; Watanabe Sokki Co. WX431). After the first cathodic sweep was imposed, subsequent cyclic sweeps gave reproducible current-potential curves (i-4 curves). These steady curves were subjected to analysis. All measurements were carried out at room temperature. All potentials in the text are presented in the normal hydrogen electrode scale. RESULTS Figure 1 shows typical steady i-4 curves under the absence (solid curve) and the presence (broken curve) of H,O,. The solid curve shows two anodic current

I

HIDEAKIKITA

peaks before oxygen evolution in the anodic sweep and two cathodic current peaks immediately after a small anodic current peak in the returning cathodic sweep, respectively. There is a double layer region at -0.3 >I#J> -1.2V.Alinearpartat4< -1.2Vgivesa slope of which value is proportional to the concentration of electrolyte in the solution. Thus, the slope is taken to represent a resistance of the solution up to the electrode surface. Comparison of the solid curve with the broken one in Fig. 1 clearly shows that the presence of H,Oz in the solution causes a shift of the curve downward at 4 > 0.2 V, but upward by an equal amount at I$ c 0.2V. Amount of the shift in both directions corresponds to a limiting diffusion current of H,O, as illustrated later. Therefore, the broken curve is explained by the superposition of the limiting oxidative or reductive diffusion current of H,O, on the solid curve. The two peaks in the anodic and cathodic sweeps are brietly discussed below. Figure 2 shows the steady i-4

I

I -1.0

I

1.0

0 $J/V

Fig. 1. Steady i+ curves at a sweep rate

I

nhe

VI

0.2 Vs- I. Full line, 1.0 M KOH; broken line, 1.0 M KOH + 0.0125 M H202.

of

-100

-Li ‘: I\

’ I

’

’

N

‘5

-50

1

’

I

I

I I

7’

4

E

;

-III

0

II

.-

50

1

Iv

,/--

v:\ .

I

I

I. 0

@

I

I

0.5

0

/

V

vs.

I

-0.5

nhe

Fig. 2. Steady i-9 curves in 1.0 M KOH at a sweep rate oTO.1V s- ‘. Reverse potentialsare 1.10 V (full line) and 0.68 V (broken line). Electric quantities in each region are I-78, Ii-36, H-2, IV-31 and V-98 mCcm_‘, respectively.

Electrochemical

behavior of H,O,

curve over a potential range from - 0.67 to 1.l V in 1.0 M KOH solution. Tilak et al.[R] observed a similar wave and ascribed the peaks to formation and reduction of silver oxides. As seen from the caption of the figure, the peak area of Ii is almost equal to that of IV, and that of I to V, respectively. Furthermore, when the sweep is reversed from 0.68 V before peak II appears, cathodic peak IV disappears as shown by the broken curve and a quantity of electricity under peak I (89 mCcm_‘) is almost equal to that of peak V (94 mCcm_‘). The amounts of electricity for each peak are not affected by stirring of solution. These facts are understood by redox reactions of surface oxides, as pointed out by Tilak et aL[8]. The standard potential of Ag/AgO is 0.342 V[9], and that of Ag,O/AgO is 0.57 V[9]. Therefore, peaks I and V are assigned as due to the reaction, Ag,0+H,0+2em and the peaks

=2Ag+20H-,

at Ag in alkaline solution

Now, ii and ii, potentialat current being zero, ie, the mixed potential +,, and reaction resistance at #I,, of solution (a+@)+, Y +,, are examined as a function composrtlon. Figure 4 shows the dependence on total H,O, concentration, [H,O,lT (sum of H,O, molecule and HO; ion concentration) of ii observed at a constant flow rate of N2 gas (cu. 600 ml min- ’ ) and at different KOH and NaClO, concentrations. ii is proportional to [H,O,], and independent of the KOH and NaClO, concentrations, ie,

(1)

II and IV are due to

2AgOtH,O+2e-eAg,O+2OH-.

(2)

Figure 3a, b shows a typical steady i-4 curve over a potential range where no surface oxides are formed. In Fig. 3a, OHconcentration is larger than that of HzO,, while the reverse is in Fig. 3b. The difference between these two figures is that in the former case the anodic limiting current density, ii, is equal to the cathodic one, iz, but in the latter ii is clearly smaller than ii. Both is and il depend on the stirring as seen

on a rotating

disc Ag electrode.

-2

-3

from fluctuations in the figures. A similar cathodic polarisation curve has been reported by Merkulova et uI.[2]

729

tog

f~~~%~~/

-1 Ml

1

Fig. 4. Hydrogen peroxide concentration dependencies of the cathodic limiting current density. AZ 0.53 M KOH, +: 0.02 M KOH + 0.06 M N&IO,.

-130 I

ii/A cn-’

= I.0 x [HIO,]f” x [K~H]~[N~C~O,]O/M.

(3)

In the absence of stirring, iz decreased to a value about one sixth of the value with a stirring. Thus, ii is taken to represent the limiting diffusion current of H,O,. Assuming a following simple equation for the diffusion, ‘d

0

0.5 cp/ V vs.

t

id = 2FD[H,O&/d

1 -0.5 nhe

’

I

1

H,O,

‘-

20 t

,

,

__%

I

1

0

0.5 Q/V

vs.nhe

Fig. 3. Steady i-4 curves at a sweep rate of 0.1 Vs-‘. (a) 0.196 M KOH +0.040 M H,O,, (b) 0.017 M KOH + 0.55 M NaCIO, + 0.025 M H,O,.

(4)

and using a value of 5 x 10m6 cm’s_‘[IO] for the diffusion coefficient, D, a thickness, d, of the diffusion layer is estimated from id as ra. lo- 3 cm which is a reasonable value. Now a question arises as to a diffusing species. H,O, is a weak acid, pK, = 11.75[11], and partly dissociates in alkaline solution as, + OH-

*HO;

+H,O.

(5)

Thus, both the concentrations of H,O, molecule and HO; ion depend largely on KOH concentration. However, as Fig. 5 shows, is is proportional to [HLOL]T and not to each concentration of H,O, molecule and HO; ion. This fact leads to the conclusion that Ihe equilibrium of reaction (5) is kept in the solution up to the electrode surface with a rate faster than the electrode reaction. Figure 5 likewise shows the dependence of il on [H,O,]T at different KOH concentrations. In the figure, ii first increases in proportion to [HIOJT and

MASATAKA

730

7

HONDA,

TAKURO

KODEPA

AND

HIDEAKI KITA

5

a \

mP .-

-2

-

! -2

-1

log Frg.

Fig. 5. Hydrogen peroxide concentration dependence of the anodic limiting current density. 0 : 0.002 M KOH + 0.20 M NaClO,, AI 0.020 M KOH + 0.060 M NaCIO,, +: 0.050 M KOH+O.O6OM NaCIO,, o: 0.10M KOH+0.060 M NaCIO,.

is equal PAM,

to limiting value when ii, but reaches exceeds a given KOH concentration. Furthermore, ifs limiting value is proportional to KOH concentration. These results are expressed as (N, gas flow rate is kept constant),

=

1.0 x [HzO,]+”

[KOHla

[KoH]/M)

+:

with

(M) at [H,O,].,.

> [KOH].

time. Similarly, a polished virgin electrode reveals the identical value with 4, immediately after the immersion into the solution and then shows a gradual shift in the negative direction with time. Dependence of the reaction resistance at I$,, composition is shown m (@/a4 9, = @,> on solution Fig. 8. The reaction resistance was obtained from the slope of the linear part near 4, in steady i-4 curve (Fig. 3a, b) by subtracting the solution resistance. Figure 8 shows that the reaction resistance is propor-

= 0.126-0.060iog[KOH]1~o x [H,O,]~[NaCIO,]”

(M-l)

\

8’

{

concentration dependence of 6,. H,O, -t NaClO,, o: with H,O*.

[NaC&,]’

1.0 x [H,O1]~[KOH]‘~o[NaC1O,]O

The effect of stirring is similarly observed as is. It is therefore concluded that the anodic limiting current density is the one either of H,02 (and HO; in equilibrium) or of OH- depending on the conditions. Next, the mixed potential 4, was obtained from Fig. 3a, b as a function of the H,O, and KOH concentrations. Results are shown in Figs 6 and 7. Figure 6 clearly shows that 4, is independent of [H~O~]T and [NaCIO,] at CH,O,]T < [KOH], whereas Fig. 7 shows that r$, is proportional to log[KOH]. Each point in Fig. 7 was taken from different experiments carried out at various H,O, concentrations under given KOH concentrations. These results are expressed as, 4,,,(V)

KOH

at [H,O& < [KOH] in agreement with the relations observed by Hurlen et a!.[11 and Bianchi et al.[12]. Merkulova et 01421 have reported that the open circuit potential is independent of a rotation speed of the disc electrode but their value is smaller than ours by about 20 mV in 0.1 M KOH + 5 x 10m4 M H,O, solution. In our experiments, 4, is also independent of solution stirring. It is, however, noted that the open circuit potential immediately after switching off polarisation is equal to 4, estimated from i - 4 curve (Fig. 3a, b) and then shifts gradually to a more negative value with

to

ii/A cm-’

7.

0

(7)

t

60mV q

0.1 -

I -3

1

“a”

I -2

‘og { [W’,]T/

I

I

-1

0

-3

Ml

Fig. 6. Hydrogen peroxideconcentration dependence of 4,. A: 0.010 M KOH, +: 0.010 M KOH + 1.0 M NaCIO,, x 0.10 M KOH, 0 : 1.0 M KOH.

-2

-1

log i W&‘zl T/~)

:

Fig. 8. H,O, concentration dependence of the reaction resistance at q5 = #J,. 0 : 0.98 M KOH,+: 0.017 M KOH + 1.0 M NaCIO,, x : 0.104 M KOH + 0.060 M NaCIOI.

Electrochemical

behavior of H,O, at Ag in alkaline solution

tiona1 to the reciprocal of [H,OJT at [HzOJT < [KOH], while independent of bath KOH and NaCIO, concentrations. These results are expressed as (0cm”)

= 0.057 [HzOz]l

*.O

9 =& x [KOH]“[NaC1O,]O

(M-l).

(8)

DISCUSSION We first assume the following mechanism to illustrate the experimental results. Equilibrium of the following acid dissociation of H,Oz holds throughout the diffusion layer in spite of the presence of the electrode reaction involving H,O,, H,O,+OH-

SHO;

+H,O.

(5)

The ionic species HO; takes part in the anodic reaction and molecular H20, takes part in the cathodic reaction by the following mechanisms: Anodic:

HO;+& HO,(a) + eHO,(a) + OH-

Cathodic:

HO;

+OH-

H,O,

+e-&

OH(a) +e-

-O,+H,O+ZeOH(a) + OH+ OH-

H,02+2e-

Pa)

--f 0, + Hz0 + e- (9b) (10) (1 la) (lib)

+20H-

(12)

where + designates a rate-determining step of the surface reaction and (a) an adsorbed state, respectively. The above mechanisms are written down so as to explain the experimental facts that the anodic limiting current reflects the diffusion of either H,O, and hence HO; at [H,O& < [OH-], equation (6a), or OHat [H,O,lT > [OH-], equation (6b) and that the cathodic limiting current reflects the diffusion of H,02. equation (3). The main purpose of the present discussion is to derive a mechanism for the surface reaction from the experimental parameters observed for the electrochemical system which does not reveal any Tafel-like polarisation curves in the anodic and cathodic branches. Reduction of 0, formed by (9b) is neglected in the present study because the rate of 0, reduction is so small compared with the rate of HzO, reduction[2].

731

calculate i, from (13) since the reaction resistance is given by (8). For example, the reaction resistance at 0.1 M H,O, is calculated from (8) as 0.55 Qcm’, and when substituted into (13), i, becomes 0.045 Acmm2. On the other hand, is at this condition is calculated from (3) as 0.1 Acm-‘. Therefore, the ratio, &,/is, becomes 0.45 and hence we cannot neglect the concentration polarisation. Littauer and Tsai[4] have also proposed the existence of concentration polarisation from their chronopotentiometric studies for the present system. Therefore, the rate equation is expressed under the condition of [H20LJT < [KOH] by taking into account the mass transfer, as follows, i, = 2Fk,[H,O,]

{ 1 - (tc+ Q/id} exp (- aF4/RT),

i, = 2Fk, [HOT]

{ 1 - (i, -t id)/&) exp (PF+/RT),

(14) (15)

where i, and i, designate cathodic and anodic current densities, k, and k, are the rate constants of (1 la) and (9a), and i, = i; = iz, respectively. The term { 1 - (i, +ia)/id) represents the degree of the H,O, concentration depression at the electrode surface. In both I, and i, of (14) and (15) we neglect respectively the reverse rate, since 4, (eg 0.126 V at 1 M KOH) is far apart from the respective equilibrium potential of the electrode reactions (10) and (12). The standard potential of reaction (IO) is - 0.076 V[9) but a real equilibrium potential must be far more negative because of an extremeIy low oxygen partial pressure in the system (N, bubbled). Equilibrium potential of reaction (12) locates near its standard value of 0.94 V. Thus at 4,. i, equals i,, and (14) and (15) give

1 = expI (a +B)FhJRT,‘.

k,lIH,OJl~,[HO~ Then, one obtains 4,

WW,l~,K,[OH-I),

= CRTI(a+B)F]

(16)

where K, designates the acid dissociation constant of HrOz, which is expressed as follows by using ion product, K,, for water K, = CH + 1 [HOi

l/CWAl = Kw[HO;]/[OH-] [H&J.

Comparison one is

(17)

of (16) with (7) leads to two relations: 2.303RT/(a

the

+ B)F = 0.060,

therefore, (i) Concentration

polarisation

cr+p=

Before we develop the rate expressions of the surface reactions, the concentration polarisation due to the mass transfer is first examined. We have presented previously[ 13, 143 the expression for the reaction resistance at $, by neglecting the reverse reactions of the respective reduction and oxidation reactions, as

*

W/W,

= 9m

= RT/Fi,

(a + fi),

and the second

is

[2.303RT/(a

+/I)F]

Using T= 293 K, = 10LL4, we have kc/k,

(13)

where i, designates the unidirectional current density at I$,, and a and fldesignate cathodic and anodic Tafel constants for the respective reduction and oxidation reactions. One can easily show that a + p = 1 on the bases of the present mechanism (9) and (11) and the experimental equation (7) [see (IS)]. Hence, we can

(ii)

Reuction

1,

log (k,K,/k,K,)

(18) = 0.126.

K, = 10-‘1~75[1 I] = 3.8 x 104.

and

(19)

K, (20)

resisrance

Next, we discuss the reaction density, i, is given as i = &-ii,.

resistance.

Net current (21)

732

MASATAKA

Cancellation of i,, i, and [HO, ] from four equations of (14), (15), (17) and (21) gives 2F{I - (k,~,lk&) exp (- ctF@/RT)} + (W/i,)

i = - (l/k,[H,O,] The above equation 4, of (16) as

is further

reduced

AND

HONDA,TAKUROKODERA

HIDMKI

By comparing each term of both (28) and (29), a and k, are determined. These calculated results are shown in

[OH-1 ext?(FdIRVI {l

+

(k,K,/k,K,)

[OH-]

(aF+/RT)/k,[Hz02]

+ (2F/i&

[exp {F(+

exp (Fe/RT)}

’

Table 1. The cathodic Tafel constant, LY,is 0.72 independent of the solution composition and is different

by introducing

2Flexp{F~4-4,)IRT)-ll

i = exp

KITA

&,,)/RT}

(22)

+ 11.

I

Differentiation (W/W,

of (22) yields at # = I#+,,

= 4, = (W/F)

{ (2/id) + (1/2Fk,

x exp W&,/W Concentration of (5) as

H202

of molecular

[H&l)

I.

(23)

Table

Substitutions of (24) and the experimentally mined equation (3) for id into (23) give = .+m=

(RT/F[H,O,]T)

(2 +x)>

[OH-]

1. Tafel constant a and the rate constant cathodic reaction in various solutions

deter(25)

1 exp W’A,lRT)lWk,.

1

0.090

0.0094

0.060

: 4

0.10 0.090 0.10

0.0245 0.0130 0.025 0.024 0.0096

0.066 0.054 0.062 0.079 - 0.032

0.196 1.0

(26)

In order to determine the value of X we have to know both the values of the rateconstant k, and the cathodic Tafel constant a. Equation (22) is rewritten as follows # = (RT/crF) In {2Fk,[H101] where

Yis defined

I+ (RT/aF)

In x(28)

as

Y= [exp{F(b-&,)/RT) - Cexp IF(d

- l]/i

-&VW

0.72

0.052

0.72 0.72 0.72 0.72 Mean value

0.024 0.035 0.023 0.074 0.039 0.041

Table 2 shows the OH- concentration dependence of the reaction resistance calculated by (25) and (26), which is in agreement with the observed one.

+ Ilhd

and can be calculated from the experimental values of &, and id as a function of 4 and i. The plot of 4 against In Y according to (28) will give a, p and k,. Figure 9 shows the results at various solution compositions. The linear line in the figure is expressed as 4 = const. + 0.035 In E

(29)

Table 2. Dependence of the reaction resistance ion concernration

P&l X .k, X RW + W

0.3

on hydroxyl

0.01

0.03

0.1

0.3

1.U

0.0129

0.0134 0.33

0.017 0.41

0.022 0.53

0.030 0.74

0.058

0.062

0.064

0.069

0.3 I

o.057

F

(iii) i, and the catalytic

0.2

decomposition

of H,O,

As already observed by several authors[3,4] and also in our experiments, 0, evolves from Ag electrode surface on open circuit. According to the present mechanism, i, = i, z i, at $ = @,, and hence i,, is given from (14), (3) and (24) as

> \ e-

k, OF the

(24)

where X is given by X = { I+ W,IK,)

(Ed= 1.0) value of ’ sp ’ and (20).

is given on the basis

IH,O,l = CHA’ZITII~+KIKJCOH-I}.

W/W,

from the reported values by Hurlen et ~I.[11 and by Merkulova et a/_[23 (a = 0.5). Mean the rate constant k, is estimated as 0.041 Mm from hence k, becomes 1.1 x 10e6 M-‘sW1

01

idjim = 2+X. 0 0

2

L

In

6

Y

Fig. 9. Relation between potential and In Yof (28). x : 0.19 M KOH+0.024MH,0~,A:0.10MKOH+0.013MH202,0: 0.10 M KOH + 0.0096 H,O,.

(301

If the catalytic decomposition of HzOl takes place electrachemically by the reactions ( 10) and (12), the rate, L,, is given as ‘CA1=

2i,

=

~ i

2

x+2

i, >

Electrochemical at 4 = $,,

behavior of H,O, at Ag in alkaline solution

rate-determining steps is equal to l/Z. However, the equation for thereaction resistanceat 4 = #J,, becomes

and 2 x+2

( >

id > i,, > 0

(32) w/w,

at 4 # 4,. These results are in agreement with the previous measurements where i,,, is the first order with respect to H,O,[63 and in maximum[2j at the open circuit potential. Therefore, we conclude that the catalytic decomposition of HzOz takes place electrochemically. (iv) Other mechanisms

examined

We discuss the following

also explain

i,-behaviour. HO;

733

two mechanisms which can The first one is

--) HO,(a)

I +m =

RT 2F [HzOJr

(* + x)’

This value is just the half of equation (25). Thus, we conclude that these alternative mechanisms are not valid and that the electrochemical reactions of H,O1 on Ag in alkaline solutions proceed through the mechanisms of (9) and (11). The above analysis will be applied without any further assumptions to more common corrosion systems. Onecan analyse the kinetic parameters of surface reactions from iz, is, 4, and the reaction resistance at

4 = 4,.

+e-,

HO,(a)+OH-fttO,+H,O+e~, for the anodic

reaction

H,Oa

REFERENCES

and

+ e-k

OH-

+ OH(a),

OH(a) + e- --) OH-, for the cathodic is

reaction,

HO; for the anodic

respectively.

The second one

+OH~,+O,+H,O+2e~ reaction H,O,

and + 2e-A,

20H-

for the cathodic reaction. Rate equations for the respective mechanisms appear the same and written in analogy with (14) and (15) as follows i, = 2Fk,[HO;]

[OH-][

1 - y]

exp@‘F4/RT)

and i, = ZFk,[H,O,]

[l -@$+xp(-or’F+,RT),

where 01’and 8’ are the corresponding Tafel constants. The observed results for the concentration dependences of 4, are explained if a’ + /3’ = 2 which appears reasonable when the symmetry factor of the respective

1. T. P. Hurlen, Y. L. Sandier and E. A. Pantler, Eleerrachim. Acra 11, 1463 (1966). 2. N. D. Merkulova, G. V. Zhutaeva, N. A. Shumilova and V. S. Bagotzky, Electrorhim. Arru 18, 169 (1973). 3. C. Iwakura, Y. Matsuda and H. Tamura, Elecrrochim. Acal 16. 471 (1971). 4. E. L. Littauer and K. C. Tsai, J. elrctrochem. Ser. 126, 1924 (1979). 5. K. Gossner and H. Bischof, 2. phys. Chem.. New Folge ‘%, 277 (1972). Lan Phuong, J. 6. M. Brezina, J. Korita and Pham-Thi electroanal. Chem. 40, 107 (1972). 7. W. Vielstich, 2. phys. Chem., New Folge 15,409 (1958). 8. B. V. Tilak, R. S. Perkins, H. A. Kozlowska and B. E. Conway, Electrochim. Acta 17. 1447 (1972). revised 2nd edn, (edited by Japan Chem. 9. Kwokubinran, Sot.) p. 1205. Maruzen, Tokyo (1975). 10. E. L. Littauer and K. C. Tsai. Eleetrochim. Acta 24,681 (1979). 11. Kogokubinran, revised 2nd edn, (edited by Japan Chem. Sot.) p. 994. Maruzen, Tokyo (1975). F. Mazza and T. Muzzini, 12. G. Bianchi, G. Caprioglio, Electrmhm. Actu 4, 232 (1961). 13. T. Kodera. H. Kita and M. Honda. Efecrrochim. Acta 17. 1361 (197i). 14. H. Kita, T. Kodera and M. Honda, J. Research Inst. Catalysts,

Hokkaido

Univ.

23, 176 (1975)