I. A. AMMARAND S. RIAD
660
Vol. 62
POLARIZATION AND OVERPOTENTIAL OF Au-CU ALLOY IN AQUEOUS SOLUTIONS BY I. A. AMMAR AND S. RIAD Department of Chemistry,Faculty of Science, University of Caiio, Cairo, Egypt Received November 29, 1957
Cathodic and anodic polarization measurements have been carried out on an alloy of Au and Cu in aqueous solutions. The overpotential for hydrogen evolution has been studied in solutions of HC1 and NaOH a t different temperatures. Attempts have been made to explain the results obtained and these have been compared with the results on pure Au and Cu. Anodic polarization has been carried out in alkaline solutions in N2 and 02 atmospheres and the time-potential relation, at a constant current density, has been established in each case. Tafel lines for oxygen evolution have been measured by different techniques and the results have been compared.
Introduction The effect of various alloying elements on hydrogen overpotential on Fe cathodes in acid solutions was studied by Stern.’ Although it is evident from his work that the Tafel line slope is sensitive to variations in the composition of the alloy, yet the theoretical interpretation for this behavior cannot be attempted easily until more experimental results are available. The aim of the present investigation is, therefore, to report the results of hydrogen overpotential measurements on alloys other than those previously studied. Au-Cu alloy was chosen because of the occurrence in the literature of reliable overpotential results on both Au2 and Cua cathodes. Results of oxygen overpotential measurements are usually less reproducible than those for hydrogen overpotential. This may be attributed to the different techniques used by various authors for oxygen overpotential measurement~.4-~ In the present investigation attempts are made to establish the Tafel lines for oxygen evolution on Au-Cu alloy in 1.0 N KOH using: (i) the constant surface method6-6; (ii) the technique of Bockris and Huq.’ In this manner the results of the two techniques could be compared. Experimental Hydrogen overpotential was measured in aqueous HC1 and NaOH solutions. HC1 solutions were repared from the constant boiling acid (prepared from “.!nalarJ’ grade HCl) with appropriate dilutions with conductivity water (K = 5 X 10-7 ohm-’ cm.-1).* NaOH was purified by crystallization, of a concentrated solution of “Analar” grade NaOH in conductivity water, under an atmosphere of pure hydrogen.9 Both HCl and NaOH solutions were pre-electrolyzed at 10-2 amp./cm.2 for 40 hours.9 The alloy electrode was in the form of a pure wire, 1 mm. diameter. The composition of the alloy was 99.5% Au and 0:5% Cu.’O The electrode was prepared by the bulb technique11 and the alloy was spot-welded to Pt; the metal to (1) M. Stern, J . Eleelrochem. Soc., 102, 663 (1955). (2) N. Pentland, J. O’M. Bockris and E. Sheldon, ibid., 104, 182 (1957). (3) J. O’M. Bockria and N. Pentland, Trans. F4TUdUU SOC.,48, 833 (1952). (4) A. Hickling and 8. Hill, Disc. Faraday Soc., 1, 236 (1947). (6) P. Jones, R. Lind and W. Wynne-Jones, Trans. Faraday SOC., 60,972 (1954). (6) P . Jones, H. Thirsk and W . Wynne-Jones, ibid., 62, 1003 (1956). . SOC.(London), (7) J. O’M. Bockris and A. Shamsul Huq., P T O CRoy. AZST, 277 (1956). (8) I. A. Ammar and 8. Awad, THIBJOURNAL, 60, 837 (1956). (9) J. O’M. Bookris and E. C. Potter, J . Chem. Phys., 20, 614 (1952). (10) Prepared by Johnson and Matthey Ltd., London. (11) J. O’M. Bockris find B. Conway, J . Sci. Inst.. 26, 283 (1948).
glass seal being made over the alloy-Pt junction. Exce t when otherwise stated this method was used throughout tEe present investigation. The overpotential cell was made of arsenic-free glass and was thermostated to &0.5O with an air thermostat. For each concentration and temperature studied several measurements were carried out using a fresh electrode and a fresh solution in each case, the results being reproducible to f10 mv. a t high current densities and f 5 mv. at low current densities. The direct method and the rapid technique were employed.’* A platinized platinum electrode in the same solution as the test cathode was used as a reference. For the work at low current densities, the current was calculated from the p.d. measured across a standard megohm resistance included in the polarizing circuit. For the measurements of oxygen overpotential using the constant surface method,‘ the alloy electrode was anodically polarized at a constant current density in 1.0 N KOH solution previously purified by cathodic pre-electrolysis. The potential, measured against a saturated calomel electrode, was recorded as a function of time until the potential was practically independent of time. This was attained when variations not exceeding 2 mv. min.-l were observed. The Tafel line was then traced rapidly to low current densities. Anodic polarization was again started a t another current density using a fresh electrode and a fresh solution and the same procedure was followed. Cathodic pre-electrolysis of the solution was carried out in a separate cell. The purified solution was then transferred with a stream of oxygen (or nitrogen) to the overpotential cell. The gas was left to bubble in the solution in the cell for several hours before the alloy electrode was lowered into the solution and the measurements were started. Oxygen overpotential was also measured using the technique described by Bockris and Huq for platinum.’ The alloy electrode was, therefore, cleaned with “Analar” chromic-sulfuric, boiling HC1 and HNOa acids. It was then washed with conductivity water and was fixed in a furnace, kept at 500°, through which a stream of oxygen was allowed to pass. Oxidation was continued for two hours after which the electrode was introduced into the previously cleaned overpotential cell. Cathodic preelectrolysis of 1.0 N KOH solution was carried out, in a separate cell, at lo-* amp./cm.2 for about 30 hours. The solution was transferred with oxygen pressure to the overpotential cell and anodic pre-electrolysis was started on an auxiliary alloy electrode with a stream of oxygen bubbling in the solution. This was carried out a t a current density of amp./cm.* for 60 hours, after which the pre-electrolysis electrode was drawn out of the solution, the test anode inserted, and the overpotential measurements were started. Current densities were calculated using the apparent surface areas.
Results and Discussion I. Hydrogen Overpotentia1.-Hydrogen overpotential was measured in the temperature range 20 to 50”. For each concentration and temperature studied more than six Tafel lines were measured in separate experiments and the mean line was computed. The results in NaOH solutions (12) A. Azzam, J. O’M. Bockris, B. Conway and H. Rosenberg, Trans. Faraday SOC.,46, 918 (1950).
POLARIZATION AND OVERPOTENTIAL OF GOLD-COPPER ALLOY
June, 1958
66 1
clearly indicate the occurrence of two Tafel line slopes in' the linear logarithmic section, ie., between 3.2 X and lo-' amp./cmS2. Below 10-7 amp./cm.2 the overpotential decreases asymptotically to the potential of the reversible hydrogen electrode. The mean results in 0.1 N NaOH at 20, 30, 38 and 48" are shown in Fig. 1 together with the corresponding relations between and the current density i for very low values of 7. It is clear from Fig. 1 that the current density corresponding to the inflection in the Tafel line increases with increase of temperature. Furthermore the relation between q and i, for very low values of q, is linear as is theoretically expected. It was, Fig. 1.-Hydrogen overpotential on Au-Cu alloy in 0.1 N NaOH: I, 20"; 11, 30"; 111, 38"; IV, 48". however, thought that the lower Tafel line slope might have been due to Pt which was used for sealing the alloy electrode to glass. Measurements were carried out, therefore, on an electrode prepared by spot-welding two wires of Pt and the AuCu alloy, sealing the Pt end to glass using the bulb technique previously mentioned and leaving part of the Pt wire free to be exposed to the solution. The mean results of five separate experiments, for electrodes having the surface area ratio (Pt/alloy) of "0.5, in 0.1 N NaOH at 20" are shown in Fig. 2. The results are reproducible to =kt0 mv. It is clear from Fig. 2 that the results cannot easily be Fig. 2.-Hydrogen overpotential on alloy-Pt cathode in 0.1 N NaOH a t 20". fitted with Tafel line slopes over any considerable range of current densities. A comparison between Figs. 1 and 2 reveals the fact that the lower Tafel line slope (Fig. 1) is an inherent property of the overpotential on the alloy electrode used rather than being due to the contamination by Pt. The same behavior as that shown in Fig. 1 was observed for alloy electrodes sealed in air, but thoroughly rubbed with fine emery paper and washed with conductivity water. In this case, two slopes were also observed in the linear logarithmic section in NaOH solutions and the values of the slopes agreed, within experimental error, with those observed for electrodes prepared by the bulb technique. Fig. 3.-Hydrogen overpotential on Au-Cu alloy in 0.1 !V HCl: I, 24"; 11, 30"; 111, 40"; IV, 50". The results in HC1 solutions indicate the occurrence of only one Tafel line slope, and the linear relation between q and i, for very low values of v, is also observed. The mean results in 0.1 N HC1 at 24,30,40 and 50" are shown in Fig. 3 together with the corresponding relations between q and i. The effect of HC1 concentration on 9 also was studied by measuring Tafel lines in 1.0, 0.1 and 0.01 N at 30". The mean results are shown in Fig. 4 from which it is clear that n decreases numerically with increase of HC1 concentration. The logarithm of the exchange current, for the mean Tafel lines, is linearly l*P related to the reciprocal of temperature. Thus for Fig. 4.-Effect of HCl concentration on hydrogen over0.1 N HC1 the relation is expressed by potential on Au-Cu alloy a t 30": I, 0.01 N ; 11, 0.10 N ; (01
r.d
d
Cd
log io = 4.3
- 3.3 x
103
(3
For 0.1 N NaOH, the equations for the high and the low current density ranges are, respectively, given by log io = -1.7 X lo3 ( l / T )
and log io = 2.5,- 3.0 X 108
(9
111, 1.0 N .
The heat of activation for hydrogen evolution was calculated from the slopes of the above relation^.^ A value of 15.2 kcal. mole-1 was obtained for 0.1 N HCI. The values calculated for 0.1 N NaOH were 13.8 and 7.8 kcal. mole-' for the lower and higher linear sections of the Tafel line. The electron number X13 was calculated from the slope of the lin(13) J. O'M. Bockria and E. C. Potter, J . Electrochem. SOC.,99, 169 (1952).
I. A. AMMAR AND 5. RIAD
662
Vol. G2
TABLE I HCI Concn., N
Temp. (“C.)
24 30 40 50 23 30
0.1
1.0
b (mv.)
92 93 93 88 90 92
io
x
(amp./cm.2)
( 1 . 8 0 . 1 ) x 10-7 (2.8 .2) x 10-7 (5.6 .3) x 10-7 ( 1 . 3 & .2) X (7.1 . 2 ) x 10-7 ( 1 . 8 f .1) x 10-6
f7
*
f3 f5 f2 f2 f5
1.6 f0 . 1 l . G f .1 1.7 f . 2 1.8 f . 1 1.8 f .2 l . G f .1
NaOH” Concn.
0.05
a
Te,mp..,
C.
bi (mv.)
20 53 2 55 f 3 30 53 f 2 38 48 52 1 4 0.10 20 52 f 2 30 53 f 3 38 55 f 3 48 50 f 2 The subscripts 1 and 2 are for the
(id1 (amp./cm.Z)
b2
(mv.)
( 5 . 9 0 . 4 ) x 10-9 122 f 3 (1.0 f 0 . 1 ) x (2.3 f . l ) X 10” 115 f 5 ( 1 . 4 =k . l ) X ( 5 . 0 i .3) X 10-8 104 i 4 (2.0 f .2) x ( 1 . 4 f .l)X 120 f G ( 3 . 2 f .3) X (2.0 f0 . 2 ) X 120 f 5 (1.8 f 0.3) x ( 3 . 6 i .2) x 10-8 120 i 4 ( 2 . 8 f .3) X ( 7 . 1 f .4) x IO-* 122 f G ( 4 . 0 f .4) x (1.3 f .2) X 110 f 7 ( 5 . 3 f .3) x lower and higher Tafel line sections, respectively.
ear relation between ?I and i at very small values of 9. The mean values of io, h and the Tafel slope Z, are given in Table I together with the corresponding 95% confidence limits.14 I n NaOH solutions the ower io values were used for the calculation of A. Calculation of X using the higher io values resulted in values of h much less than one in contradiction to theoretical expectations. Previous work on hydrogen overpotential on pure Au cathodes has indicated the occurrence of two Tafel line slopes (0.076 and 0.097 v.) in HC1 solutions and one slope (0.116-0.123 v.) in NaOH solutions.2 X has been given as 2 for HCI solutions (and also 0.1 N NaOH) and one for alkaline sohtions less concentrated than 0.1 N . The results on pure Au have been explained on the basis of a ratedetermining electrochemical desorption in HC1 solutions. The occurrence of two Tafel line slopes has been attributed to the dependence of the potential, 6,(at the Gouy-Helmholtz boundary) on the metal-solution p.d., A b , particularly in the vicinity of the electrocapillary ma~imum.~5I n dilute alkaline solutions less than 0.1 N ) a slow discharge mechanism has been suggested.2 However, the results of the present investigation on Au-Cu alloy show the occurrence of one Tafel slope in HCl sohtions and two slopes in NaOH solutions. Values of X (Table I) are near to two for both acid and alkaline solutions. The behavior of Tafel lines in alkaline solutions is the same for electrodes sealed: (i) under hydrogen; and (ii) in air. The above comparison clearly indicates that the Parameters of hydrogen evolution on Au cathodes are changed by alloYiW it with as little 0.5% CU. This is in agreement with the observation of Stern1 that the Tafel slope and the exchange current for hydrogen evolution on pure Fe are altered by alloying Fe lvith various elements. The fact that the Au-Cu alloy, studied in the present investigation, is characterized by a Tafel slope of 0.088-0.093 v., in (14) 0. Davies, “Statistical Methods in Research and Production,” Oliver and Boyd, London (1949). (15) J. O’M. Bockris, I. A. Ammar and A. K. Huq, THISJOURNAL, 61, 879 (1957).
c
(id2
(amp./cm.Z)
A
10-6 10-6 10-6 10-6 10-6
2.3 f0.2 2.4 f .2 2.3 f . 1 2.3 f .2 2.1 f0 . 2 2.4 f . 1 2.4 f .1 2.1 f . 1
HCI solutions, indicates that the rate of hydrogen evolution is governed either by a slow proton discharge or by a slow electrochemical desorption. Since X lies between 1.6 and 1.8 (cf. Table I) the slow electrochemical desorption is indicated. This conclusion is identical with that previously drawn for pure Cu16 and Au2 cathodes in acid solutions. It is clear from Table I that, in NaOH solutions, the lower Tafel line slope lies between 0.050 and 0.055 v. Similar slopes were reported in the literature. Thus slopes of 0.053-0.054 v. were observed, in HC1 solutions, for Pt cathodes sealed in air, and were attributed to a slow surface migration of H atoms prior to evolution of Hz molecules.16 Surface migration of H atoms is expected to be slow, and hence rate determining, in presence of poisons which block the surface thus decreasing the rate of H migration. This explains why slopes of 0.053-0.054 v. were observed for Pt cathodes sealed in air and not for cathodes sealed under pure hydrogen.’6 The results given in Table I are for cathodes sealed under a pure hydrogen atmosphere in absence of surface contaminations by poisons. Hence, the lower Tafel line slopes in NaOH solutions may not be due to the slow surface migration of H atoms. The state of the metal surface plays an important role in modifying the kinetics of hydrogen evolution, Thus, Busing and E(auzniann17 have showii that the Tafel equation may not be valid for electrode reactions on non-uniform surfaces. Breaks in the Tafel lines for hydrogen evolution on electrodeposited Cu in HCI solutions have, therefore, been attributed to the heterogeneity of the surface such that hydrogen is evolved on two nlain kinds of sites of adsorption; each being effective a t a different current density range.18 Surface heterogeneity may also account for the break in the Tafel line on Au-Cu alloy in NaOH solutions. The fact that (16) B. Conway and J. O’M. Bockris, J . Chem. Phye., 26, 532 (1957). (17) W. Busing and W. Ihuzmrtnn, ibid., 20, 1129 (1952). (18) 8. Wakkad, I. A. Ammar and 1%. Sabry, J . Chem. Soc., 3020
(195G).
B
POLARIZATION AND OVERPOTENTIAL OF GOLD-COPPER ALLOY
June, 1958
only one Tafel line slope is observed in HC1 solutions in contrast t o the results in NaOH solutions (Table I) may indicate that the catalytic centers are different for the reactions taking place in acid and alkaline solutions. However, another explanation for the results of the present investigation is considered below. Values of 9 in NaOH solutions (Table I) indicate a rate-determining electrochemical desorption or catalytic combination. The catalytic combination mechanism is associated with a slope of 0.03 v., even when the electrode potential is in the vicinity of the potential of the electrocapillary maximum. l 6 The fact that slopes of ca. 0.03 v. are not observed in the results of the present investigation, together with the following discussion, favor a rate-determining electrochemical desorption for hydrogen evolution on the Au-Cu alloy in NaOH solutions. I n alkaline solutions it is possible that the discharge of hydrogen takes place from water molecules rather than from H30+ions.19 Hence, when the electrochemical desorption is rate determining, the reactions occurring at the cathode may be represented by
+ e- + M +OH- + MH, fast
(1)
where M represents the metal MH + OH- +M HzO + e-, fast and
(2)
HzO
+
MH
+ HzO + e- +Hz + M + OH-, slow
where a, is the activity of HzO molecules in the initial state, X is the fraction of the surface covered with adsorbed atomic hydrogen, 0,0' and 0"are symmetry factors. I n reaction 2 the act,ivity of OH- is much greater than that of adsorbed H atoms and it is, therefore, taken as a constant and is included in lcz (equation 5 ) . When the electrochemical desorption is rate determining the discharge step may be taken as a reversible reaction and hence VI N VZ>>, VI. The expression for X , therefore, becomes [cf. equations 4, 6 , G and ref. 131 -N a l / ( a l
+ az)
(7)
Id
.o
,do
110
ID0
140
t h'"1.
Fig. 5.-Anodic polarization of Au-Cu alloy in 1.0 N KOH at 30": E us. normal hydrogen potential; electrode area 0.63 cm.2; Oz atmosphere: I, 2.5 X amp., 11, 1.5 X amp. Nzatmosphere: 111, 1.5 X amp., IV, 5 X 10-6 amp.
xm 4
(3)
Reaction 2 represents the reverse of the discharge reaction 1. The rates associated with the above reactions, taking the, initial state in the Helmholtz double layer, may be given by
x
04,
663
110 r d
'x
Fig. B.-Tafel lines for Oz evolution after nnodic olarira tion for 3-4 hours at: I, 2.5 X amp. ( 0 2 1 , If, 1.5 10-4 amp. (02); 111, 1.5 X amp. (Nz).
01
(u
_ I
-7
101 C d
..
Fig. 7.-Tnfel lines for 0 2 evolut,ion: I-using the technique of Bockris and Huq (IO), 11, IIIk-after anodic, polarization for 24 hours a t 1.5 X amp. in Nzand OZ atmospheres, respectively.
rived by Bockris and Potter13 for acid solutions, where (Sa) applies for the low current density range and (8b) for high current densities. From (4), (5), (6) and (8a)
+ +
Equation 7 can be simplified under two limiting P' P " ) (A+ - d)F] ($-)) ) ~ - (P conditions: (i) a2 >> al and; (ii) al >> a2. VI = I C ~ ( U ~exp RT Hence If the assumption is now made that a , in the double x = (a,/az) @a) layer is equal to that in the bulk of solution, the exor pression for the net cathodic current becomes (8b) X = l ( P f P' 4- 8") (A+ - +,PI i = kjexp[RT Equations 8a and 8b are identical with those deHence, according to (10) t,he Tafel slope, a t low cur(19) R . Parsons and J. O'M. Bockris, Trans. Faraday SOC.,47, 914 rent densities, is given by (1951).
[
I. A. AMMARAND S. RIAD
G 64 nrn
At high current densities when X = 1 the slope (cf. equation 6 ) becomes (bA+/bIn i) =
- P"F(1 -RT&$/aA+)
(12)
From (11) the slope is 0.04 v. when (dd/bAd) = 0, i.e., a t potentials far from that of the electrocapillarymaximumand whenp = p' = p" = 0.5. The occurrence of slopes of 0.050-0.055 v. in the results in alkaline solutions (Table I) can be accounted for on the basis of (11) if (d~$/dAd)has a positive value, Le., when the electrode potential is near to that of the electrocapillary maximum. l5 As the electrode potential becomes more negative the fraction X + 1, the slope is then given by (12) and the electrode potential is shifted from that of the electrocapillary maximum such that (d+/dA$~) -P 0. Hence, the slope approaches 0.12 v. which is experimentally observed for the high current density range. The above argument, which could have been checked had it been possible to get some reliable information as to the potential of the electrocapillary maximum for Au-Cu alloy in NaOH solutions, provides an attempt to explain the occurrence of two slopes in alkaline solutions. 11. Oxygen Overpotential.-Anodic polarization of the Au-Cu alloy was carried out at a constant current density in 1.0 N KOH at 30". Measurements were taken in nitrogen as well as oxygen atmospheres. The results are shown in Fig. 5 from which it is clear that the potential increases rapidly at the beginning, and with increase of time it gradually approaches a steady value. It is also clear from Fig. 5 that the steps of anodic oxidation could not be detected in the results even with polarizing currents as low as 5 X 10+ amp. I n all cases studied, the steady state potential is higher than the equilibrium potentials for the three possible oxides of Au20and this steady state may, therefore, be attributed to oxidation of the metal surface as well as oxygen evolution The absence of the various steps of oxidation can be explained on the basis of the amphoteric character of the oxides of gold which dissolve in strong alkaline solutions forming aurite and aurate.20 In an oxygen at(20) S. Wakkad and A. Din, J . Chem. SOC.,3098 (1854).
Vol. 62
mosphere the steady-state potential is higher than that in a nitrogen atmosphere a t the same polarizing current density. Thus, a t 2.4 X IOm4 amp./ cm.2 this potential in an oxygen atmosphere is 1.145 v. with respect to the normal hydrogen potential, while in a nitrogen atmosphere it is 0.925 V.
For the measurements of Tafel lines, the anode was left polarized, a t a constant current density, for 3 to 4 hours and the anode potential was then rapidly traced to low current densities. The overpotential was calculated from the measured potential taking the e.m.f. of the hydrogen-oxygen cell as 1.23 v. Tafel lines thus obtained are shown in Fig. 6. It is clear from this figure that, in an oxygen atmosphere, 9 is linear to the logarithm of the current density when the initial current density used for anodic polarization is low. Tafel line slopes are in the vicinity of 0.2 v. However, at high polarizing current densities this linearity is not observed. In a nitrogen atmosphere the Tafel lines are linear, with a slope of ca. 0.1 v., down to about 1 X amp./cm.2. Below this current density a negative deviation from linearity is observed and the potential changes appreciably with time. The overpotential, a t a constant current density, is usually higher in an oxygen atmosphere than the corresponding value in a nitrogen atmosphere. Another set of Tafel lines was obtained after anodic polarization, a t a constant current density, for 24 hours. The results are shown in Fig. 7. It is clear from Fig. G and 7 that prolonged anodic polarization increases the overpotential, and this increase is greater in an oxygen atmosphere than that observed in a nitrogen atmosphere. The Tafel line slopes are not, however, affected by prolonged anodic polarization. This agrees with the observation of Jones, Lind and Wynne-Jones who find that the Tafel slope on Pb02 electrode does not change with the time of polarization.5 Figure 7 also shows an example of the Tafel lines obtained using the technique of Bockris and Huq.' The slope in this case is 0.12 v. and the overpotential is lower than the values obtained by the constant surface method. The authors wish to express their thanks to Prof. A. R. Tourky for his interest in this work.
*
..
