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The electrochemistry of silver in koh at elevated temperatures—II. Cyclic voltammetry and galvanostatic charging studies
THE ELECTROCHEMISTRY OF SILVER IN KOH AT ELEVATED TEMPERATURES--II. CYCLIC VOLTAMMETRY AND GALVANOSTATIC CHARGING STUDIES BRUCE G. POUND,* DIGBY D. MACDONALD~~ and JOHN W. TOMLINSON$ * Department of Chemistry, University of Auckland, Auckland,New Zealand t Department
of Metallurgical
Engineering, The Ohio State University, Columbus, OH 43210, U.S.A.
$ Department of Chemistry, Victoria University of Wellington, Wellington, New Zealand (Receiwd 7
August 1979)
electrochemistry of polycrystalliue silver in 1mol kg- ’ KOH at temperatures between295K and 478K hasbeenstudiedusingcyclic voltammetry and cyclical galvanostatic charging. Comparison of the observed peak potentials(cyclic voltammetry) and arrest potentials (galvanostatic charging) with calculated Abstract -The
equilibrium potentials shows that previously proposed mechanisms for the oxidation-reduction behaviour of silver in alkaline solution are viable. However, detailed analyses of the charges passed during each arrest indicate that the actual mechanisms involved are more complex thanhasheenassumedin the past. Finally the total anodic charge capacity of polycrystalline silver in 1 mol kg-‘.KOH was found to passthrough a maximum at - 370 K (- KWC), thereby indicating that
in aqueous systems at elevated temperatures is of considerable interest in a variety of fundamental and applied technology fields, including power generation and storage, extractive metallurgy, thermodynamics, materials science, and electrochemical kinetics. Because of the experimental difficulty of gaining access to the interior of a pressure vessel containing a high temperature-high pressure aqueous system, the number of electrochemical techniques that have been used for the investigation of metals in this environment is limited. Steady-state or slow-sweep potentiodynamic methods have been most commonly employed[ll5], but in some cases other relaxation techniques[6] have been used. In this paper we report a study of the electrochemical behavior of silver in 1 mol kg-’ KOH over the temperature range of 295-478 K using cyclic voltammetry and the galvanostatic charging technique. Silver was chosen because it has been studied extensively at ambient temperature and pressure[7-231 and, accordingly, the processes that occur at the surface have been well defined. Potentiostatic transient and variable frequency ac impedance studies will be described in a subsequent publication. EXPERIMENTAL Apparatus and electrode preparation The
316
SS high-pressure/high-temperature
-.--6 To whom correspondence should be addressed. Key Words: Silver, anodic oxidation, potassium droxide, elevated temperatures.
elec-
trochemical was fitted tamination
cell was of conventional with a Teflon liner to
design[4], minimize
and con-
of the hydroxide solution by steel corrosion products. A cylindrical silver working electrode (99.99% purity, 0.01 m long, 0.005 m dia, 1.57 x 10m4 m2 apparent surface area) was mounted on a steel rod that protruded through a Teflon sheath. A leak-free fit between the upper end of the working electrode and the Teflon sheath was ensured by compression nuts on the external end of the steel rod. The other end of the silver electrode was masked from the solution by a Teflon end cap. A symmetrical electric field was ensured by use of a cylindrical silver counter electrode tocated concentrically around the working electrode. Both the working and counter electrodes were cleaned prior to each run. In the case of the working electrode, the surface was polished successively with 320, 600, and 800 Sic paper. The electrode was then washed in ammonia followed by double-distilled water, degreased in acetone, washed with water again and then dried. A similar procedure was adopted for preparing the counter electrode, except that a less stringent sequence was observed during polishing the surface. The effect of pretreatment on the surface activity of silver has been examined by Morley[Z4] using differential capacitance measurements. Morley found considerable variation in surface activity for different surface pretreatment procedures. Accordingly, the strict sequence of pretreatment steps described above was adopted in this work in an attempt to eliminate spurious effects associated with surface preparation. Reference electrode
An external calomel reference electrode of the type described previously[4] was used for all potential
hy563
BRUCE G. POUND,
564
DIGBYD.
MACDONALD
measurements. However, it was desirable to relate the measured electrode potentials to a meaningful thermodynamic scale. Because the potential of an external non-pressure balanced electrode contains contributions from two irreversible processes (thermal diffusion due to the temperature gradient, and a streaming potential due to the pressure gradient), it was necessary to calibrate the reference electrode system against a thermodynamically well-defined internal electrode. Following Macdonald[25], the external reference electrode is conveniently represented by the following thermal cell
: IEexll
SCE (298 K)
SHE (298 K)
for which the measured potential is given by E en =Er+EdP-0.098V
(4)
The SHEs are included in the analysis of cell (3) for conceptual reasons only; they were not used in the actual experimental measurements. Expanding EAT in terms of the thermal coefficients for the electrodes[25] gives E,,, = -00.98 + [’
(g)
dT + E,J
dT
(5)
Isoth.Hg-HgO
KOH
SHE (-U
WE U)
(1)
where EAT is the potential of the standard hydrogen electrode thermocell containing a 1 mol kg- ’ KOH bridge, and Er is the potential of the working electrode with respect to the SHE at the same temperature. It is this latter quantity that ls compared directly with calculated equilibrium potentials. The observed potential, E,,, is related to the quantities defined above by E cx*=E’r+Edl.+EEU
- 0.244V
(2)
where EM is the isothermal liquid junction potential between the see and the KOH solution at 298 K. E, , was estimated using the Hendersen equation[24] anld was found to be about 0.009 V. The streaming potential contribution to EAT was determined by locating identical Hg/HgO electrodes inside and outside the pressure vessel. The vessel was then pressurized with nitrogen at 295 K to 0.4 MPa to force KOH solution through the asbestos plug. No significant potential difference between the two Hg-HgO electrodes was observed. In a second experiment, the nitrogen pressure in the cell was increased in steps to 0.8 MPa after obtaining a steady-state cyclic voltammogram for silver at 0.4 MPa. At each pressure, cyclic voltammograms for silver were recorded using both the internal and external Hg-HgO reference electrodes. No appreciable shifts in the voltammograms were produced by the pressure increments, nor were the curves found to have moved relative to that for ambient pressure. Accordingly, we conclude that the streaming potential contribution E,,, is negligible. The thermal liquid junction potential contribution to E,, was evaluated experimentally using the thermocell
r--Ecn-----l SHE Hg/HgO
AND JOHN W. TOMUNSON
KOH
SHE
HgO/Hg
were (LYE/aT&,,, and (BE/dT)Ilo,hH _HgO are the thermal coefficients and isothermal toe tpi. cients for the SHE and mercury-mercuric oxide electrodes, respectively. Substitution of values for these coefficients listed by deBethune et a/[271 and assuming that they are temperature independent therefore yields E,,, = E7., - 0.249 x lo-’
(T - 298)
(6)
where ETJ is the thermal liquid junction potential for the KOH bridge. Experimentally determined values for E,, at 348 K and 388 K were found to be 0.0115 k 0.001 V and 0.0255 + 0.001 V, respectively. Suhstitution of these values into equation (6) therefore yields values of 0.024 V and 0.048 V for ErJ at 348 K and 388 K, respectively. deBethune et aI[27] recommended that (aE/aT),, = 0.0005 V/K and, by making use of the observation of Seys and Van Haute[28] that for many systems this coefficient is not highly temperature dependent, we calculate EfJ to be 0.025 Vat 348 K and 0.045 V at 388 K. These .-raluesagree sufficiently well with the experimentally determined values given above that the coefficient (aE/JT),, sx 0.0005 V/K probably yields a fair estimate for the thermal liquid junction potential for 1 mol kg-r KOH. The thermal coefficient for the SHE given by deBethune et a1[27) has been shown recently by Macdonald[25] to be approximate and is probably reliable only for temperatures lower than about 393 K. At higher temperatures, (dE/dT),,, becomes increasingly nonlinear with temperature, and the value for this coefficient assumed above[27] tends to overestimate the thermal potential for the SHE. Accordingly, the more precise values of Macdonald[25] were used in correcting the measured potential of the working electrode with respect to the external SCE cf (2). E’, = Em - Erh,sm - 0.5 x lo-‘(T
- 298) - 0.235
(7)
The thermal potential of the SHE is taken as 0,0.030, O.OSX,0.085, 0.098, 0.100, and 0.096 at 298 K, 333 K, 373 K, 423 K, 473 K, 523 K, and 573 K, respectively, and values for intermediate temperatures were obtained by interpolation. Cyclic voitammetry Cyclic voltammograms of silver in 1 mol kg-’ KOH at elevated temperatures were obtained using a Chemical Electronics Type 4O-3A potentiostat which was driven by a Philips PM 5168 function
c3)
The electrochemistryof silver in KOH at elevated temperatures
tials for peaks Al, AZ, and C2 (Fig. 1) still occurred after eight cycles, altbough the shift per cycle was greatly diminished. The general shape of the voltammogram at 295 K is typical of those obtained by previous workers[7, IS] under ambient conditions. Three principal anodic (oxidation) peaks (A,, A,, A3) are observed prior to oxygen evolution at all temperatures employed. However, at 348 K peak A, is resolved into two subpeaks ; this behavior was not observed at any of the other temperatures studied in this work. Tilak et aI[15] have observed a fourth anodic peak at temperatures below ambient (243 K). However, as the temperature is raised, the oxygen evolution current in this potential region becomes appreciable, so that at room temperature the fourth anodic peak cannot be resolved. The additional anodic peak detected by Tilak et al does not appear to be equivalent to either of the subpeaks A, or Aj observed in this work at 348 K.
generator. The cyclic voltammograms were recorded using an X-Y recorder. Galuanostatic
charging
studies
Chronopotentiograms were recorded using a potentiostat operating in the constant current mode. The charging curves were recorded on a fast response X-t recorder. RESULTS Cyclic ooltammetry (1) Genewl. Typical cyclic voltammograms for silpotassium hydroxide at six temperaverin lmolkg-’ tures between 295 and 478 K are shown in Fig. 1. These curves were obtained at a voltage sweep-rate of 16.8 mV s-l, and were found to be superimposable after four to five cycles at all temperatures except 478 K. In this latter case, changes in the peak poten80 ”
E z 40 a
0
A2 Al I
u g so
c3 295
i=IT!!
E/V
F d
160
0.4
0 E/V
-0.4
K
C2
-0.4
(SCE.
0
0.4
(SCE,
295
295
Kl
-0.4 K)
-0.4
0
E/V
ISCE,
0 E/V
(SCE, 295
0.4 295 K)
0.4 KI
80
-0.4 E/V
0 (SCE,
0.4 295
0.6 K)
565
E/V
(SCE,
295
KJ
Fig. 1. Cyclic voltammogramsfor silver in 1 mol kg- 1 KOH as a function of temperature. Voltage sweep rate = 16.8mVs-l.
566
BRUCE G. POVND, DIGBY D. MACDONALD
Two principal cathodic (reduction) peaks (C,, C,) were observed in the present study on sweeping the potential in the negative direction from the oxygen evolution region. An additional cathodic peak (C,) was observed at 388 K (Fig. 1 (c)) between C, and C,. However, the peak height of C, is very much less than that for either of the two principal cathodic peaks, so that it probably arises from a minor charge transfer process at the interface. Consequently, this peak will not be discussed further in this paper. Increasing temperature is seen to have a marked effect on the shapes of the voltammograms, and also on the relative heights of the peaks. For example, peak A, grows with respect to peak A2 with increasing temperature over the range 295 K to 458 K, but this trend is reversed on increasing the temperature to 478 K. Furthermore, the cathodic peak C, is not observed at this highest temperature, even though its conjugate anodic peak (A3) is. Also, a small anodic peak, 11 occurs at intermediate temperatures on the cathodic sweep immediately before Cs. Although this peak is not observed at the lower temperatures, it becomes quite distinct at 388 and 428 K. However, when the temperature is raised to 458 K, only a dip in the curve is evident, while at 478 K peak II has disappeared completely. In a room temperature study of thin films of silver, Stonehart and Portante[9] observed a similar peak which they associated with the completion of Ag,O, formation. At all temperatures, cyclic sweeping caused a general increase in peak currents and also an increase in the charge transferred at each peak. The first oxidation peak (Al), in particular, was observed to become more prominent on cycling. The observed increase in current on continued cycling is probably due to progressive roughening of the surface.
AND JOHN W. TOMLINSON
T < 428 K). At the two highest temperatures (458 K and 478 K), the observed uncertainty in peak potential (f60 mV) precluded any quantitative correlation of the type given in Fig. 2. Peak current data are plotted in Figs. 3 to 5 as a function of the square root of sweeprate[6] at various temperatures between 295 and 428 K ; at higher temperatures lack of reproducibility again precluded quantitative correlation. In general, linear plots of i, 11s ei” are observed. However, these correlations do not extrapolate- to the origin, as predicted by simple diffusion theory under quiescent conditions[6] or expected if the peak arises from a single charge transfer process[ 6,293. The $,vs vil’ behaviour observed in this work at 295 K is m agreement with that observed previously[15] at ambient temperature for the same range of sweep-rates. Gaivanostatic
charging
Galvanostatic charging curves for silver in 1 mol kg-’ KOH over the temperature range 295-478 K are shown in Fig. 6. These curves were obtained using a current density of 5.7 mA cm-’ (apparent surface area basis), and were found to be superimposable after 3 to 5 cycles at all temperatures except 478 K (cf cyclic voltammetry). The form of the curve at 295 K is typical of that observed in past studies of silver at a similar temperature[l4]. The potentials of the arrests are generally in good agreement with previously measured values. At all temperatures, the potential initially changes rapidly with time in the noble direction upon anodic charging due principally to double layer charging. At longer times, potential arrests are observed on both the oxidation and reduction cycles. Except for temperatures of 428 K and 478 K, the potential changes very (2) Sweep-rate dependence. Cyclic voltammograms rapidly between each arrest thereby indicating were obtained at a number of voltage sweep rates over effective inhibition of the charge transfer processes by the range 9 to 40mV s-’ at each temperature. The the oxidation products formed at the surface. peak currents, and in many cases the peak potentials, Generally good agreement is obtained between the were found to be sensitive to sweep rate, as has been galvanostatic charging results and the cyclic voltamobserved for other metals in alkaline solutions at mograms, at least as far as the number of processes elevated temperatures[2,4]. The peak currents in- indicated by each technique is concerned. Thus, for crease with sweep-rate for all peaks over the range of temperatures below 478 K. three charge transfer protemperature considered, whereas the effect of sweepcesses on anodic charging are indicated, whereas two rate on the peak potential varied with the particular principal arrests are observed on cathodic charging peak and the temperature. from the oxygen evolution region. However, at 478 K, On the anodic sweep, the peak potential for Al was only two anodic and one cathodic arrests are shown by found to be independent of sweep-rate at all temperathe charging curves, whereas at least three anodic tures. By contrast A3 (and also A3’ at 348 K) shift to reactions were detected by cyclic voltammetry (Fig. I). more positive potentials with increasing sweep-rate at However, the anodic charging curve at this temperaeach temperature. Peak A2 was intermediate in beture exhibits a gradual change in potential with time havior, exhibiting a transition from a dependence on subsequent to arrest EA2, so that any additional sweep rate up to 348 K to a nondependence at 388 K processes might not be expected to exhibit clearly and above. On the cathodic sweep, C,, and also Cd, defined arrests. shift to more cathodic potentials with increasing An important feature of the charging curves shown sweep-rate. However, C, exhibits no apparent sweepin Fig. 6 is the considerable transient overpotentials at rate dependence except at 388 K. short times that are indicated for the formulation and Peak potential versus logarithm of sweep rate, In(v), reduction of various surface phases. Thus, at 295 K the data are plotted in Fig. 2 for temperatures up to 428 K. apparent transient overpotential associated with arAlthough the range of sweep-rate considered is limited, rest EA3 is about 0.1 V. The size of the transient it is apparent that linear correlations are obtained at overpotential is found to be dependent on temperaall temperatures except at 388 K. In this case, the ture ; in general it decreases as the temperature is raised observed nonlinearity is barely significant above the as expected from the effect of temperature on the rates experimental uncertainty in E, ( f 15 mV for 295 K -ZZ of electrochemical processes.
The electrochemistry
of silver in KON
at elevated temperatures
567
620 A3
3oa
25C 2.8
3.5
360
320
300
2.1
Fig. 2. Dependence
2.8 I”(“/mvs-’
3.5
J
with thermodynamic
behnuior
Previous cyclic voltammetry work under ambient conditions by Stonehart and coworkers[7-91 using sweep reversal techniques has indicated that peaks A2 and C, are conjugate, as are peaks A, and C,. Thus, reversal of the active to noble sweep at potentials between peaks A, and A, results only in cathodic peak Cs. However, reversal of the sweep at potentiafs on the noble side of peak A, gives rise to both reduction peaks C, and C,, except at 478 K in which case peak C, is not observed even if the forward sweep is carried well into oxygen evolution (Fig. l(f)). No clearly de&ted reduction
peak conjugate
2.8 In WmVs-’
3.5
1
of peak potential (E,) on In(v) for peaks A2, A3, C2, C3 and C4 as a function of temperature.
DISCUSSION
Comparison
It
2.1
to Al is apparently
observed.
It is now generally accepted[7, 11. 161 that the first oxidation peak Al (Fig. 1) and arrest Al’ (Fig. 6) on anodic charging are due to the dissolution of silver to form AgO-. This hypothesis is in keeping with the predicted thermodynamic behavior[30], which indicates that in alkaline solution dissolution should take place at a potential that is more active than that for the formation of the lowest oxide. Cyclic vohammetric peaks A2 and C2, and potential arrests A2’ and C’2’ observed on galvanostatic charging, have been attributed to the formation and reduction of Ag,O. For example, Briggs et al.[20], using X-ray diiraction techniques, have identified this phase as being formed on silver during anodic charging through arrest A2’. Similarly, peaks A3 and C3, and arrests A3’ and C3’
BRUCE G. POUND,DIGBY D.MACDONAL.D
568
ANDJOHN
W.
TOMLINS~N
4
2
0
6
"llz/,~t/2,-1/2
Fig. 5. Dependenceofpeak current($1on u”’ for peak C2 as a functionof temperature.
have been attributed to the formation and reduction of Ag,Oa [Ag(I) A&III) O,] according to the reactionC9, 11, 151 Ag,O + 20H-~2 Ag,O,
Fig. 3. Dependence of peak current (i,) on VI’* for peaksAl and A2 as a function of temperature. 300
I (~1
I
I
I
A3
200
0 120
-4
I
I
I
I
I
I
/
,/
I 388
K
60 ‘: ‘m .40 I
0
0
2
6 yv2,m”&2
Fig. 4. Dependence of peak current (i,) on LV for peaks A3 and C3 as a function of temperature.
(8)
Although a quantitative comparison between kinetic peak and arrest potentials, and thermodynamic equilibrium potentials[30] is not possible, the general consistency of the redox mechanism discussed above can be tested by taking into account the expected signs for the overpotentials contained in the kinetic parameters. Thus, since the overpotentials for anodic (oxidation) and cathodic (reduction) processes are expected to be positive and negative, respectiveIy, the peak and arrest potentials (EAx, E.--) for conjugate phenomena will be related to the equilibrium potential (E,,,,) for the appropriate reaction according to the following inequality E,,
100
+ I-I,0 + 2e-
2 Ecq.x z &X
(9)
Peak and arrest potentials are compared with the appropriate equilibrium potentials in Table 1. In general, inequality (9) is satisfied by the data listed in Table 1, particularly in view of the estimated uncertainty in the peak and arrest potentials of +0.05 V. Accordingly, the assignments of A2(A2’) and C2(C2’) to the formation and reduction of Ag,O and of A3(A3’) and C3(C3’) to the couple Ag,O/Ag,O, are consistent with the variable temperature studies reported here. Charge considerations and mechanism Cyclic voltammetry and cyclic galvanostatic charging are convenient and effective techniques for evaluating the charge capacities of electrodes. These two experimental techniques have also been used extensively for the evaluation of kinetic parameters for both film formation/reduction phenomena and redox reactionsC6, 7-9, 291. In this section, the cyclic voltammetry and cyclic galvanostatic charging results are used to examine the effect of temperature on the charge capacity of polycrystalline silver in 1 mol kg-’ potassium hydroxide solution.
The electrochemistry
I
-0.2
I
of silver in KOH at devated temperatures
I
388 K
I
569
I
0 2
0.2 0.4 0.6 a
-0.6
100
470
200
300
0.6
0
100
200
K
-0.4 > -0.2 L’d 2 0 E’A2
E’c2
0.2 0.4
E’A3
0
Fig. 6. Galvanostatic
50
loo f/S
150
charging curves for silver in 1 mol kg-’ KOH as a functionof temperature. Charging current = 5.7 mA cm-‘. Voltage USsee at 22°C.
Typical charge versus temperature profiles observed on cyclic galvanostatic charging between the limits of hydrogen evolution and oxygen evolution are summarized in Fig. 7. These data were calculated from the durations of the arrests shown in Fig. 6 and the applied current density. The total anodic and cathodic charges are taken as the sums of the contributions observed on anodization and cathodization, respectively. No direct measurement of the contributions to the charges plotted in Fig. 7 from dissolution processes were made. The total charges plotted in Fig. 7 range from 0.3 C cm- ’ to almost 1.0C cm-+. Since a monolayer of electrochemically adsorbed oxygen ions is equivalent to -0.4mC crne2, the observed anodic charges indicate the formation of oxide films ranging in thickness from - 0.08 pm to -0.25 pm (based on the apparent
area). The film thickness data obtained in this study at 295 K are typical of those that have been observed in previous studies[20] of the anodic oxidation of silver in hydroxide solution under ambient conditions. An interesting and important feature of the data plotted in Fig. 7 is that the total anodic and cathodic charges observed on anodization and subsequent cathodization pass through maxima as a function of temperature. In the present study, the optimum temperature for the formation of surface oxides on silver in 1 mol kg- ’ KOH is found to be - 370 K ( - loO°C) ; at higher temperatures the total charge that can be stored at the interface as oxidation products decreases sharply with increasing T. The contribution to the total anodic charge from the surface
570
BRUCEG. POUND,DIGBY D. MACDDNALDAND JOHNW. TOMLNS~N
Table 1. Comparison of peak (cyclic voltammetry) and arrest (galvanostatic potentials with calculatedequilibrium poteutials[30] for reactionsinvolvingsilver kg-’ KOH at elevated temperatures
charging) in 1 mol
VK 295
348
388
428
458
478
0.41
0.28 0.25 -0.03
0.23 0.19 -0.15
0.14
0.22
0.33 0.31 0.07
-0.23
0.03 0.05 -0.28
0.60 0.51 0.35 0.27 0.30
0.42 0.34 0.27 0.18 0.25
0.35 0.28 0.21 0.14 0.21
0.26 0.20 0.13 0.11 0.15
0.17 0.07 -0.08 -
0.14 0.05 0.02 -0.21 0.06
0.76 0.69 0.61 0.56 0.56
0.74 0.60 0.54 0.48 0.49
0.66 0.55 0.48 0.46 0.48
0.59 0.50 0.42 0.42 0.43
0.53 0.36 0.28
0.59
Ex = potential for peak X as observed using cyclic voltammetry, arrest Y as observed on galvanostatic chargbtg.
0.31
E; = potential for
1.0
0.8
0.6
0.4
N 0.2 L V V 0 :
n
0
C3’
2 3 -0.2 C2’ -0.4
k
\ \ /
-0.6
/ \ \
/
/
4/ \
//
-0.8
-1.0
300
400
5aa
T/K Fig. 7. Charge UStemperature
profiles for arrests A2’, A3’, IX, C3’ for polycrystalline KOH. QA = total anodic charge. Qc = total cathodic charge. Charging current
silver in 1 mol kg-’ = 5.7 mA cm- *.
The electrochemistryof silver in KOH at elevated temperatures
Additional evidence for the validity of this viewpoint is provided by considering the ratio of charges associated with the various potential arrests observed on cyclical galvanostatic charging. Thus, as shown in Fig. 8, the ratio of the total anodic to total cathodic charge increases with temperature from 1.0 at 295 K to -2.5 at 423 K. This increase may reflect the fact that cathodization does not reduce all of the oxide film formed during prior anodization ; a more likely explanation is that increasing temperature favors parallel dissolution processes, in which case some of the oxidation products are not localized at the interface and are therefore not quantitatively available for reduction on cathodization. This view is supported by the cyclic voltammograms plotted in Fig. 1 which show that the dissolution peak Al in general becomes more prominent at higher temperatures (except at 478 K). While the ratio of the total anodic to total cathodic charge increases only slightly with increasing temperature, the ratios for the conjugate charge transfer processes exhibit much larger variations. For example, Q,,,,/QcJ, increases by a factor of two upon increasing the temperature from 295 K to 428 K, whereas Q,+2,/QczJ passes through a shallow minimum over the same temperature range. Furthermore, the charge ratios for neighbouring arrests Qaa,fQn2. and also exhibit strong variations with temperaQw/Qc,~ ture as shown by the data plotted in Fig. 8. It is difficult to reconcile these observed charge ratios with a simple mechanism that involves the following reactions
oxidation of silver to Ag,O [peaks A2 (Fig. l), arrest A2’ IFis. 611is found to be minor and indenendent of ten&&u~, except at temperatures abovk -46OK (Fig. 7). At these higher temperatures, QA2, increases sharply with increasing T, such that at 478 K it becomes the dominant contribution to the total anodic charge. However, over most of the temperature range of interest, the total charge capacity of silver in 1 mol kg-l KOH is determined by the third principal oxidation process; that is, by those processes that give rise to peak A3 (cyclic voltammetry, Fig. 1) and potential arrest A3’ (galvanostatic charging, Fig. 6). The fact that the QAzz and Q,,, us temperature profiles are not complimentary suggests that the third anodic arrest cannot be attributed soley to the oxidation of Ag,O to AgO, as has been frequently assumed in the past. This point is discussed in greater detail below. In contrast to the different temperature profiles for QA2, and Q,,,, noted above, those for the reduction processes C2’ and C3’ are similar, in that both pass through maxima at -370 K (- 1OO’C). Furthermore, Qcl, is greater than Qc-*, which is contrary to the behaviour exhibited by the conjugate arrests A2’ and A3’ (Qazs < Qaj, T < 460 K). These relationships provide additional evidence that the potential arrests A2’, C2’ and A3’, C3’ do not arise from simple conjugate charge transfer reactions as implied by most mechanisms that have been proposed to account for the electrochemical behaviour of the silver electrode in alkaline solution.
,
I
571
I
QA/%
A3’/A2’
_
Fig. 8. Charge
ratio
us tcmpcrature
profilm
for various
arrests
obserwd
during
galvanostatic
silver in 1 mol kg-’ KOH. Charging current = 5.7mA cm-*.
charging
of
572
BRUCE G. POUND, DIGBY D.
Ag + 20H2AG + 20H-,
AgzO
+
201~
-
*’Ai’ AgOAZ.AL?
A c2.c2,
-G A3~y
+ H,O f e-
Ag,O + H,O + 2eAg,O,
MACDONALD
(10) (11)
+ Hz0 + 2e- (12)
Thus, if the arrests A2’, C2’, and A3’, C3’ arise from conjugate processes, as described above, then QA3./QA2. and Q&Q,--, should both be equal to unity, assuming complete oxidation of Ag,O to Agz02, in contrast to the complex variations with temperature for both quantities shown in Fig. 8. It can be argued that oxidation of Ag to Ag,O, or even directly to Ag,Oz, can occur in the region of potential arrest A3’, might be expected to be greater than so that QAY/QAZ’ one. This is indeed the case for temperatures less than WK, but it is not so for the highest temperature (478 K). Also, no obvious explanation is apparent for the maximum in this ratio at 370 K. Likewise, Ag,O, could continue to be reduced in the region of potential arrest C2, thereby giving rise to a ratio Qc2’/Qc3. that is greater than one, as observed. Miller[19] found in his rotating ring-disk electrode study under ambient conditions that on reduction of Ag,O, the surface is almost immediately covered with Ag,O. His results indicate a dissolution-precipitation mechanism for this process of the type A&O, 2AgO-
+ 2e- + 2AgO-
(dissolution)
+ H,O * Ag,O + 20H-
(13)
(precipitation)(l4)
Since not all of the dissolved intermediate may pra cipitate out as Ag,O, the net effect is that the charge due to reduction ofAg,O should be less than indicated by that for the reduction ofAg,O, ; that is, Qcz,/Qcs, is expected to be less than unity rather than in excess of one as observed. Accordingly, the dissolutionprecipitation mechanism proposed by Miller[l9] does not appear to account for the variable temperature results report in this study. Finally, the corn$ lexity of the oxidation-reduction mechanism for silver is shown by the variable sweep rate cyclic voltammetry data plotted in Figs. 2 to 5. Thus, the fact that some of the E, us In(u) and i, 0s vl” correlations exhibit nonlinearity, and particularly that the i, us ul” plots in general do not pass through the origin, indicates that the peaks observed by cyclic voltammetry do not arise from simple charge transfer reactions at least over the range of conditions employed in this work. Clearly, a much more detailed study of the kinetics of the electrochemical oxidation and reduction of silver and its oxides is required if the mechanisms for these important processes are to be delineated. SUMMARY AND CONCLUSIONS
The electrochemical behaviour of polycrystalline silver in 1 mol kg-l KOH over the temperature range 295 K to 478 K has been investigated using cyclic voltammetry and cyclic galvanostatic charging. The findings of this study are summarized as foIlows : (1) The processes that are generally assumed to occur
AND
JOHN W. TOMLWSCIN
on the anodic oxidation of silver (Ag + AgO-, Ag -+ Ag,O, Ag,O -, Ag202) are consistent with the calculated thermodynamic equilibrium behaviour of silver in 1 mol kg- ’ KOH over the temperature range of interest. (2) Detailed analysis of the charge associated with each charge transfer process indicates that the processes that give rise to the potential arresta A2, C2 and A3, C3 are not strictly conjugate; that is the reduction does not take place via the simple reversal of the forward oxidation step. The complexity of the redox mechanism is indicated by the nonlinearity of peak potential us In(u) and peak current us ti”’ plots do not pass through the origin. (3) The total anodic charge capacity of polycrystalline silver in 1 mol kg-’ KOH was found to pass through a maximum at - 370 K (- 100°C), thereby indicating that the performance of a polycrystalline silver anode is optimal at this temperature.
REFERENCES
1. D. D. Macdonald, in Modern Aspecrs ofElecrrochemisrry (Edited \- ~~- bv -. B. E. Conwav N.<. and J. O’M. Bockris) Vol. il. -.D.
141, Pklnml Press. (1975). 2. D. D. Macdonald, J. elecrrochem. Sot. 121, 651 (1974). 3. D. D. Mardonald; Electrochim. Acto 21, 169 (1976). 4. D. D. Macdonald and D. Owen, J. electrochem. Sot. 12% 317 (1973). 5. D. D. Macdonald and D. Owen, paper presented at the International Conference on High Temperature High
Pressure
(NA.C.E.1.
Electrochemistry University
in
Aqueous
of Surrey (1973).
Solutions
6.
istry Plenum Pre& N.Y. (1977). . 7. P. Stonehart, Elecrrochim. Acta 13, 1989 (1968). 8. P. Stonehart, If. A. Kozlowska and B. E. Conway. Proc. R. Sot.. Land. A310, 541 (1969). 9. P. Stonehart and F. P. Portante, Electrochim. Acto 13, 1805 (1968). 10. R. Luther and F. Pokorney, Z. Anorg. U. Allgem. Chem. 57, 290 (1908). 11. T. P. Dirkse and D. B. De Vries, 1. phys. Chem 43, 107 (1959). 2, 117 (1972). 12. N. A. Hampson, Elecwochemisrry 13. N. A. Hampson. J. B. Lee and J. R. Morley, Elecrrochim. Acta 16, 637 (1971). 14. A. Hickling and D. Taylor, Disc. Faraday Sot. I, 277 (1947). 15. B. V. Tilak, R. S. Perkins, H. A. Kozlowska and 8. E. Conway, EIectrochim. Acta 17, 1447 (1972). 16. R. D. Giles and J. A. Harrison, J. Electroanul. Chem. 27, 161 (1970). and H. R. Thirsk, J. 17. R. D. Gjles. J. A. Harrison elecrroanui. Chem. 22., 375 (19691 . 18. M. Fleischrnann. D. J. Lax and H. R. Thirsk, Trans. Faraday Sot. 64, 3137 (1967). 19. B. Mill&, J. electrochem. Sot. 117, 491 (1970). 20. G. W. D. Briees. M. Fleischmann. D. J. Lax and H. R. Tbirsk, Trans~F~raday Sot. 64. 3i20 (1968). 21. M. Fleischmann, D. J. Lax and H. R. Thirsk, Truns. Faraday Sot. 64, 3128 (1968). 22. T. G. Clarke, N. A. Hampson, J. B. Lee, J. R. Morley and B. Scanlon, Con. J. Chem. 46, 3437 (1968). 23. B. E. Conway, M. A. Sattar and D. Gilroy, Elecrrochim. Acra 14, 677 (1969). Thesis, Loughborough, University of 24. J. R. Marty, Technology (1969).
The electrochemistry
of silver in KOH at elevated temperatures
25. D. D. Macdonald, Corrosion 34, 75 (1978). 26. P. Henderson, Z. phys. Chem. 59, 118 (1907); 63, 325 (1908). 27. A. 3. deBethune, T. S. Licht and N. Swendeman, J. ekctrochem. Sot. 106. 617 (1959). 28. A. A. Scys and A. A. Van Haute, paper presented to rhe International Conference on High Temperature High
573
Pressure Electrochemistry in Aqueous Solutions (N.A.C.E.), University of Surrey (1973). 29. S. Srinivasan and E. Gileadi, Electrochim. Actu 11, 321 (1966). 30. B. G. Pound, D. D. Macdonald and J. W. Tomlinson, Electrochim. Acta 24, 929 (1979).
of Metallurgical
Engineering, The Ohio State University, Columbus, OH 43210, U.S.A.
$ Department of Chemistry, Victoria University of Wellington, Wellington, New Zealand (Receiwd 7
August 1979)
electrochemistry of polycrystalliue silver in 1mol kg- ’ KOH at temperatures between295K and 478K hasbeenstudiedusingcyclic voltammetry and cyclical galvanostatic charging. Comparison of the observed peak potentials(cyclic voltammetry) and arrest potentials (galvanostatic charging) with calculated Abstract -The
equilibrium potentials shows that previously proposed mechanisms for the oxidation-reduction behaviour of silver in alkaline solution are viable. However, detailed analyses of the charges passed during each arrest indicate that the actual mechanisms involved are more complex thanhasheenassumedin the past. Finally the total anodic charge capacity of polycrystalline silver in 1 mol kg-‘.KOH was found to passthrough a maximum at - 370 K (- KWC), thereby indicating that
in aqueous systems at elevated temperatures is of considerable interest in a variety of fundamental and applied technology fields, including power generation and storage, extractive metallurgy, thermodynamics, materials science, and electrochemical kinetics. Because of the experimental difficulty of gaining access to the interior of a pressure vessel containing a high temperature-high pressure aqueous system, the number of electrochemical techniques that have been used for the investigation of metals in this environment is limited. Steady-state or slow-sweep potentiodynamic methods have been most commonly employed[ll5], but in some cases other relaxation techniques[6] have been used. In this paper we report a study of the electrochemical behavior of silver in 1 mol kg-’ KOH over the temperature range of 295-478 K using cyclic voltammetry and the galvanostatic charging technique. Silver was chosen because it has been studied extensively at ambient temperature and pressure[7-231 and, accordingly, the processes that occur at the surface have been well defined. Potentiostatic transient and variable frequency ac impedance studies will be described in a subsequent publication. EXPERIMENTAL Apparatus and electrode preparation The
316
SS high-pressure/high-temperature
-.--6 To whom correspondence should be addressed. Key Words: Silver, anodic oxidation, potassium droxide, elevated temperatures.
elec-
trochemical was fitted tamination
cell was of conventional with a Teflon liner to
design[4], minimize
and con-
of the hydroxide solution by steel corrosion products. A cylindrical silver working electrode (99.99% purity, 0.01 m long, 0.005 m dia, 1.57 x 10m4 m2 apparent surface area) was mounted on a steel rod that protruded through a Teflon sheath. A leak-free fit between the upper end of the working electrode and the Teflon sheath was ensured by compression nuts on the external end of the steel rod. The other end of the silver electrode was masked from the solution by a Teflon end cap. A symmetrical electric field was ensured by use of a cylindrical silver counter electrode tocated concentrically around the working electrode. Both the working and counter electrodes were cleaned prior to each run. In the case of the working electrode, the surface was polished successively with 320, 600, and 800 Sic paper. The electrode was then washed in ammonia followed by double-distilled water, degreased in acetone, washed with water again and then dried. A similar procedure was adopted for preparing the counter electrode, except that a less stringent sequence was observed during polishing the surface. The effect of pretreatment on the surface activity of silver has been examined by Morley[Z4] using differential capacitance measurements. Morley found considerable variation in surface activity for different surface pretreatment procedures. Accordingly, the strict sequence of pretreatment steps described above was adopted in this work in an attempt to eliminate spurious effects associated with surface preparation. Reference electrode
An external calomel reference electrode of the type described previously[4] was used for all potential
hy563
BRUCE G. POUND,
564
DIGBYD.
MACDONALD
measurements. However, it was desirable to relate the measured electrode potentials to a meaningful thermodynamic scale. Because the potential of an external non-pressure balanced electrode contains contributions from two irreversible processes (thermal diffusion due to the temperature gradient, and a streaming potential due to the pressure gradient), it was necessary to calibrate the reference electrode system against a thermodynamically well-defined internal electrode. Following Macdonald[25], the external reference electrode is conveniently represented by the following thermal cell
: IEexll
SCE (298 K)
SHE (298 K)
for which the measured potential is given by E en =Er+EdP-0.098V
(4)
The SHEs are included in the analysis of cell (3) for conceptual reasons only; they were not used in the actual experimental measurements. Expanding EAT in terms of the thermal coefficients for the electrodes[25] gives E,,, = -00.98 + [’
(g)
dT + E,J
dT
(5)
Isoth.Hg-HgO
KOH
SHE (-U
WE U)
(1)
where EAT is the potential of the standard hydrogen electrode thermocell containing a 1 mol kg- ’ KOH bridge, and Er is the potential of the working electrode with respect to the SHE at the same temperature. It is this latter quantity that ls compared directly with calculated equilibrium potentials. The observed potential, E,,, is related to the quantities defined above by E cx*=E’r+Edl.+EEU
- 0.244V
(2)
where EM is the isothermal liquid junction potential between the see and the KOH solution at 298 K. E, , was estimated using the Hendersen equation[24] anld was found to be about 0.009 V. The streaming potential contribution to EAT was determined by locating identical Hg/HgO electrodes inside and outside the pressure vessel. The vessel was then pressurized with nitrogen at 295 K to 0.4 MPa to force KOH solution through the asbestos plug. No significant potential difference between the two Hg-HgO electrodes was observed. In a second experiment, the nitrogen pressure in the cell was increased in steps to 0.8 MPa after obtaining a steady-state cyclic voltammogram for silver at 0.4 MPa. At each pressure, cyclic voltammograms for silver were recorded using both the internal and external Hg-HgO reference electrodes. No appreciable shifts in the voltammograms were produced by the pressure increments, nor were the curves found to have moved relative to that for ambient pressure. Accordingly, we conclude that the streaming potential contribution E,,, is negligible. The thermal liquid junction potential contribution to E,, was evaluated experimentally using the thermocell
r--Ecn-----l SHE Hg/HgO
AND JOHN W. TOMUNSON
KOH
SHE
HgO/Hg
were (LYE/aT&,,, and (BE/dT)Ilo,hH _HgO are the thermal coefficients and isothermal toe tpi. cients for the SHE and mercury-mercuric oxide electrodes, respectively. Substitution of values for these coefficients listed by deBethune et a/[271 and assuming that they are temperature independent therefore yields E,,, = E7., - 0.249 x lo-’
(T - 298)
(6)
where ETJ is the thermal liquid junction potential for the KOH bridge. Experimentally determined values for E,, at 348 K and 388 K were found to be 0.0115 k 0.001 V and 0.0255 + 0.001 V, respectively. Suhstitution of these values into equation (6) therefore yields values of 0.024 V and 0.048 V for ErJ at 348 K and 388 K, respectively. deBethune et aI[27] recommended that (aE/aT),, = 0.0005 V/K and, by making use of the observation of Seys and Van Haute[28] that for many systems this coefficient is not highly temperature dependent, we calculate EfJ to be 0.025 Vat 348 K and 0.045 V at 388 K. These .-raluesagree sufficiently well with the experimentally determined values given above that the coefficient (aE/JT),, sx 0.0005 V/K probably yields a fair estimate for the thermal liquid junction potential for 1 mol kg-r KOH. The thermal coefficient for the SHE given by deBethune et a1[27) has been shown recently by Macdonald[25] to be approximate and is probably reliable only for temperatures lower than about 393 K. At higher temperatures, (dE/dT),,, becomes increasingly nonlinear with temperature, and the value for this coefficient assumed above[27] tends to overestimate the thermal potential for the SHE. Accordingly, the more precise values of Macdonald[25] were used in correcting the measured potential of the working electrode with respect to the external SCE cf (2). E’, = Em - Erh,sm - 0.5 x lo-‘(T
- 298) - 0.235
(7)
The thermal potential of the SHE is taken as 0,0.030, O.OSX,0.085, 0.098, 0.100, and 0.096 at 298 K, 333 K, 373 K, 423 K, 473 K, 523 K, and 573 K, respectively, and values for intermediate temperatures were obtained by interpolation. Cyclic voitammetry Cyclic voltammograms of silver in 1 mol kg-’ KOH at elevated temperatures were obtained using a Chemical Electronics Type 4O-3A potentiostat which was driven by a Philips PM 5168 function
c3)
The electrochemistryof silver in KOH at elevated temperatures
tials for peaks Al, AZ, and C2 (Fig. 1) still occurred after eight cycles, altbough the shift per cycle was greatly diminished. The general shape of the voltammogram at 295 K is typical of those obtained by previous workers[7, IS] under ambient conditions. Three principal anodic (oxidation) peaks (A,, A,, A3) are observed prior to oxygen evolution at all temperatures employed. However, at 348 K peak A, is resolved into two subpeaks ; this behavior was not observed at any of the other temperatures studied in this work. Tilak et aI[15] have observed a fourth anodic peak at temperatures below ambient (243 K). However, as the temperature is raised, the oxygen evolution current in this potential region becomes appreciable, so that at room temperature the fourth anodic peak cannot be resolved. The additional anodic peak detected by Tilak et al does not appear to be equivalent to either of the subpeaks A, or Aj observed in this work at 348 K.
generator. The cyclic voltammograms were recorded using an X-Y recorder. Galuanostatic
charging
studies
Chronopotentiograms were recorded using a potentiostat operating in the constant current mode. The charging curves were recorded on a fast response X-t recorder. RESULTS Cyclic ooltammetry (1) Genewl. Typical cyclic voltammograms for silpotassium hydroxide at six temperaverin lmolkg-’ tures between 295 and 478 K are shown in Fig. 1. These curves were obtained at a voltage sweep-rate of 16.8 mV s-l, and were found to be superimposable after four to five cycles at all temperatures except 478 K. In this latter case, changes in the peak poten80 ”
E z 40 a
0
A2 Al I
u g so
c3 295
i=IT!!
E/V
F d
160
0.4
0 E/V
-0.4
K
C2
-0.4
(SCE.
0
0.4
(SCE,
295
295
Kl
-0.4 K)
-0.4
0
E/V
ISCE,
0 E/V
(SCE, 295
0.4 295 K)
0.4 KI
80
-0.4 E/V
0 (SCE,
0.4 295
0.6 K)
565
E/V
(SCE,
295
KJ
Fig. 1. Cyclic voltammogramsfor silver in 1 mol kg- 1 KOH as a function of temperature. Voltage sweep rate = 16.8mVs-l.
566
BRUCE G. POVND, DIGBY D. MACDONALD
Two principal cathodic (reduction) peaks (C,, C,) were observed in the present study on sweeping the potential in the negative direction from the oxygen evolution region. An additional cathodic peak (C,) was observed at 388 K (Fig. 1 (c)) between C, and C,. However, the peak height of C, is very much less than that for either of the two principal cathodic peaks, so that it probably arises from a minor charge transfer process at the interface. Consequently, this peak will not be discussed further in this paper. Increasing temperature is seen to have a marked effect on the shapes of the voltammograms, and also on the relative heights of the peaks. For example, peak A, grows with respect to peak A2 with increasing temperature over the range 295 K to 458 K, but this trend is reversed on increasing the temperature to 478 K. Furthermore, the cathodic peak C, is not observed at this highest temperature, even though its conjugate anodic peak (A3) is. Also, a small anodic peak, 11 occurs at intermediate temperatures on the cathodic sweep immediately before Cs. Although this peak is not observed at the lower temperatures, it becomes quite distinct at 388 and 428 K. However, when the temperature is raised to 458 K, only a dip in the curve is evident, while at 478 K peak II has disappeared completely. In a room temperature study of thin films of silver, Stonehart and Portante[9] observed a similar peak which they associated with the completion of Ag,O, formation. At all temperatures, cyclic sweeping caused a general increase in peak currents and also an increase in the charge transferred at each peak. The first oxidation peak (Al), in particular, was observed to become more prominent on cycling. The observed increase in current on continued cycling is probably due to progressive roughening of the surface.
AND JOHN W. TOMLINSON
T < 428 K). At the two highest temperatures (458 K and 478 K), the observed uncertainty in peak potential (f60 mV) precluded any quantitative correlation of the type given in Fig. 2. Peak current data are plotted in Figs. 3 to 5 as a function of the square root of sweeprate[6] at various temperatures between 295 and 428 K ; at higher temperatures lack of reproducibility again precluded quantitative correlation. In general, linear plots of i, 11s ei” are observed. However, these correlations do not extrapolate- to the origin, as predicted by simple diffusion theory under quiescent conditions[6] or expected if the peak arises from a single charge transfer process[ 6,293. The $,vs vil’ behaviour observed in this work at 295 K is m agreement with that observed previously[15] at ambient temperature for the same range of sweep-rates. Gaivanostatic
charging
Galvanostatic charging curves for silver in 1 mol kg-’ KOH over the temperature range 295-478 K are shown in Fig. 6. These curves were obtained using a current density of 5.7 mA cm-’ (apparent surface area basis), and were found to be superimposable after 3 to 5 cycles at all temperatures except 478 K (cf cyclic voltammetry). The form of the curve at 295 K is typical of that observed in past studies of silver at a similar temperature[l4]. The potentials of the arrests are generally in good agreement with previously measured values. At all temperatures, the potential initially changes rapidly with time in the noble direction upon anodic charging due principally to double layer charging. At longer times, potential arrests are observed on both the oxidation and reduction cycles. Except for temperatures of 428 K and 478 K, the potential changes very (2) Sweep-rate dependence. Cyclic voltammograms rapidly between each arrest thereby indicating were obtained at a number of voltage sweep rates over effective inhibition of the charge transfer processes by the range 9 to 40mV s-’ at each temperature. The the oxidation products formed at the surface. peak currents, and in many cases the peak potentials, Generally good agreement is obtained between the were found to be sensitive to sweep rate, as has been galvanostatic charging results and the cyclic voltamobserved for other metals in alkaline solutions at mograms, at least as far as the number of processes elevated temperatures[2,4]. The peak currents in- indicated by each technique is concerned. Thus, for crease with sweep-rate for all peaks over the range of temperatures below 478 K. three charge transfer protemperature considered, whereas the effect of sweepcesses on anodic charging are indicated, whereas two rate on the peak potential varied with the particular principal arrests are observed on cathodic charging peak and the temperature. from the oxygen evolution region. However, at 478 K, On the anodic sweep, the peak potential for Al was only two anodic and one cathodic arrests are shown by found to be independent of sweep-rate at all temperathe charging curves, whereas at least three anodic tures. By contrast A3 (and also A3’ at 348 K) shift to reactions were detected by cyclic voltammetry (Fig. I). more positive potentials with increasing sweep-rate at However, the anodic charging curve at this temperaeach temperature. Peak A2 was intermediate in beture exhibits a gradual change in potential with time havior, exhibiting a transition from a dependence on subsequent to arrest EA2, so that any additional sweep rate up to 348 K to a nondependence at 388 K processes might not be expected to exhibit clearly and above. On the cathodic sweep, C,, and also Cd, defined arrests. shift to more cathodic potentials with increasing An important feature of the charging curves shown sweep-rate. However, C, exhibits no apparent sweepin Fig. 6 is the considerable transient overpotentials at rate dependence except at 388 K. short times that are indicated for the formulation and Peak potential versus logarithm of sweep rate, In(v), reduction of various surface phases. Thus, at 295 K the data are plotted in Fig. 2 for temperatures up to 428 K. apparent transient overpotential associated with arAlthough the range of sweep-rate considered is limited, rest EA3 is about 0.1 V. The size of the transient it is apparent that linear correlations are obtained at overpotential is found to be dependent on temperaall temperatures except at 388 K. In this case, the ture ; in general it decreases as the temperature is raised observed nonlinearity is barely significant above the as expected from the effect of temperature on the rates experimental uncertainty in E, ( f 15 mV for 295 K -ZZ of electrochemical processes.
The electrochemistry
of silver in KON
at elevated temperatures
567
620 A3
3oa
25C 2.8
3.5
360
320
300
2.1
Fig. 2. Dependence
2.8 I”(“/mvs-’
3.5
J
with thermodynamic
behnuior
Previous cyclic voltammetry work under ambient conditions by Stonehart and coworkers[7-91 using sweep reversal techniques has indicated that peaks A2 and C, are conjugate, as are peaks A, and C,. Thus, reversal of the active to noble sweep at potentials between peaks A, and A, results only in cathodic peak Cs. However, reversal of the sweep at potentiafs on the noble side of peak A, gives rise to both reduction peaks C, and C,, except at 478 K in which case peak C, is not observed even if the forward sweep is carried well into oxygen evolution (Fig. l(f)). No clearly de&ted reduction
peak conjugate
2.8 In WmVs-’
3.5
1
of peak potential (E,) on In(v) for peaks A2, A3, C2, C3 and C4 as a function of temperature.
DISCUSSION
Comparison
It
2.1
to Al is apparently
observed.
It is now generally accepted[7, 11. 161 that the first oxidation peak Al (Fig. 1) and arrest Al’ (Fig. 6) on anodic charging are due to the dissolution of silver to form AgO-. This hypothesis is in keeping with the predicted thermodynamic behavior[30], which indicates that in alkaline solution dissolution should take place at a potential that is more active than that for the formation of the lowest oxide. Cyclic vohammetric peaks A2 and C2, and potential arrests A2’ and C’2’ observed on galvanostatic charging, have been attributed to the formation and reduction of Ag,O. For example, Briggs et al.[20], using X-ray diiraction techniques, have identified this phase as being formed on silver during anodic charging through arrest A2’. Similarly, peaks A3 and C3, and arrests A3’ and C3’
BRUCE G. POUND,DIGBY D.MACDONAL.D
568
ANDJOHN
W.
TOMLINS~N
4
2
0
6
"llz/,~t/2,-1/2
Fig. 5. Dependenceofpeak current($1on u”’ for peak C2 as a functionof temperature.
have been attributed to the formation and reduction of Ag,Oa [Ag(I) A&III) O,] according to the reactionC9, 11, 151 Ag,O + 20H-~2 Ag,O,
Fig. 3. Dependence of peak current (i,) on VI’* for peaksAl and A2 as a function of temperature. 300
I (~1
I
I
I
A3
200
0 120
-4
I
I
I
I
I
I
/
,/
I 388
K
60 ‘: ‘m .40 I
0
0
2
6 yv2,m”&2
Fig. 4. Dependence of peak current (i,) on LV for peaks A3 and C3 as a function of temperature.
(8)
Although a quantitative comparison between kinetic peak and arrest potentials, and thermodynamic equilibrium potentials[30] is not possible, the general consistency of the redox mechanism discussed above can be tested by taking into account the expected signs for the overpotentials contained in the kinetic parameters. Thus, since the overpotentials for anodic (oxidation) and cathodic (reduction) processes are expected to be positive and negative, respectiveIy, the peak and arrest potentials (EAx, E.--) for conjugate phenomena will be related to the equilibrium potential (E,,,,) for the appropriate reaction according to the following inequality E,,
100
+ I-I,0 + 2e-
2 Ecq.x z &X
(9)
Peak and arrest potentials are compared with the appropriate equilibrium potentials in Table 1. In general, inequality (9) is satisfied by the data listed in Table 1, particularly in view of the estimated uncertainty in the peak and arrest potentials of +0.05 V. Accordingly, the assignments of A2(A2’) and C2(C2’) to the formation and reduction of Ag,O and of A3(A3’) and C3(C3’) to the couple Ag,O/Ag,O, are consistent with the variable temperature studies reported here. Charge considerations and mechanism Cyclic voltammetry and cyclic galvanostatic charging are convenient and effective techniques for evaluating the charge capacities of electrodes. These two experimental techniques have also been used extensively for the evaluation of kinetic parameters for both film formation/reduction phenomena and redox reactionsC6, 7-9, 291. In this section, the cyclic voltammetry and cyclic galvanostatic charging results are used to examine the effect of temperature on the charge capacity of polycrystalline silver in 1 mol kg-’ potassium hydroxide solution.
The electrochemistry
I
-0.2
I
of silver in KOH at devated temperatures
I
388 K
I
569
I
0 2
0.2 0.4 0.6 a
-0.6
100
470
200
300
0.6
0
100
200
K
-0.4 > -0.2 L’d 2 0 E’A2
E’c2
0.2 0.4
E’A3
0
Fig. 6. Galvanostatic
50
loo f/S
150
charging curves for silver in 1 mol kg-’ KOH as a functionof temperature. Charging current = 5.7 mA cm-‘. Voltage USsee at 22°C.
Typical charge versus temperature profiles observed on cyclic galvanostatic charging between the limits of hydrogen evolution and oxygen evolution are summarized in Fig. 7. These data were calculated from the durations of the arrests shown in Fig. 6 and the applied current density. The total anodic and cathodic charges are taken as the sums of the contributions observed on anodization and cathodization, respectively. No direct measurement of the contributions to the charges plotted in Fig. 7 from dissolution processes were made. The total charges plotted in Fig. 7 range from 0.3 C cm- ’ to almost 1.0C cm-+. Since a monolayer of electrochemically adsorbed oxygen ions is equivalent to -0.4mC crne2, the observed anodic charges indicate the formation of oxide films ranging in thickness from - 0.08 pm to -0.25 pm (based on the apparent
area). The film thickness data obtained in this study at 295 K are typical of those that have been observed in previous studies[20] of the anodic oxidation of silver in hydroxide solution under ambient conditions. An interesting and important feature of the data plotted in Fig. 7 is that the total anodic and cathodic charges observed on anodization and subsequent cathodization pass through maxima as a function of temperature. In the present study, the optimum temperature for the formation of surface oxides on silver in 1 mol kg- ’ KOH is found to be - 370 K ( - loO°C) ; at higher temperatures the total charge that can be stored at the interface as oxidation products decreases sharply with increasing T. The contribution to the total anodic charge from the surface
570
BRUCEG. POUND,DIGBY D. MACDDNALDAND JOHNW. TOMLNS~N
Table 1. Comparison of peak (cyclic voltammetry) and arrest (galvanostatic potentials with calculatedequilibrium poteutials[30] for reactionsinvolvingsilver kg-’ KOH at elevated temperatures
charging) in 1 mol
VK 295
348
388
428
458
478
0.41
0.28 0.25 -0.03
0.23 0.19 -0.15
0.14
0.22
0.33 0.31 0.07
-0.23
0.03 0.05 -0.28
0.60 0.51 0.35 0.27 0.30
0.42 0.34 0.27 0.18 0.25
0.35 0.28 0.21 0.14 0.21
0.26 0.20 0.13 0.11 0.15
0.17 0.07 -0.08 -
0.14 0.05 0.02 -0.21 0.06
0.76 0.69 0.61 0.56 0.56
0.74 0.60 0.54 0.48 0.49
0.66 0.55 0.48 0.46 0.48
0.59 0.50 0.42 0.42 0.43
0.53 0.36 0.28
0.59
Ex = potential for peak X as observed using cyclic voltammetry, arrest Y as observed on galvanostatic chargbtg.
0.31
E; = potential for
1.0
0.8
0.6
0.4
N 0.2 L V V 0 :
n
0
C3’
2 3 -0.2 C2’ -0.4
k
\ \ /
-0.6
/ \ \
/
/
4/ \
//
-0.8
-1.0
300
400
5aa
T/K Fig. 7. Charge UStemperature
profiles for arrests A2’, A3’, IX, C3’ for polycrystalline KOH. QA = total anodic charge. Qc = total cathodic charge. Charging current
silver in 1 mol kg-’ = 5.7 mA cm- *.
The electrochemistryof silver in KOH at elevated temperatures
Additional evidence for the validity of this viewpoint is provided by considering the ratio of charges associated with the various potential arrests observed on cyclical galvanostatic charging. Thus, as shown in Fig. 8, the ratio of the total anodic to total cathodic charge increases with temperature from 1.0 at 295 K to -2.5 at 423 K. This increase may reflect the fact that cathodization does not reduce all of the oxide film formed during prior anodization ; a more likely explanation is that increasing temperature favors parallel dissolution processes, in which case some of the oxidation products are not localized at the interface and are therefore not quantitatively available for reduction on cathodization. This view is supported by the cyclic voltammograms plotted in Fig. 1 which show that the dissolution peak Al in general becomes more prominent at higher temperatures (except at 478 K). While the ratio of the total anodic to total cathodic charge increases only slightly with increasing temperature, the ratios for the conjugate charge transfer processes exhibit much larger variations. For example, Q,,,,/QcJ, increases by a factor of two upon increasing the temperature from 295 K to 428 K, whereas Q,+2,/QczJ passes through a shallow minimum over the same temperature range. Furthermore, the charge ratios for neighbouring arrests Qaa,fQn2. and also exhibit strong variations with temperaQw/Qc,~ ture as shown by the data plotted in Fig. 8. It is difficult to reconcile these observed charge ratios with a simple mechanism that involves the following reactions
oxidation of silver to Ag,O [peaks A2 (Fig. l), arrest A2’ IFis. 611is found to be minor and indenendent of ten&&u~, except at temperatures abovk -46OK (Fig. 7). At these higher temperatures, QA2, increases sharply with increasing T, such that at 478 K it becomes the dominant contribution to the total anodic charge. However, over most of the temperature range of interest, the total charge capacity of silver in 1 mol kg-l KOH is determined by the third principal oxidation process; that is, by those processes that give rise to peak A3 (cyclic voltammetry, Fig. 1) and potential arrest A3’ (galvanostatic charging, Fig. 6). The fact that the QAzz and Q,,, us temperature profiles are not complimentary suggests that the third anodic arrest cannot be attributed soley to the oxidation of Ag,O to AgO, as has been frequently assumed in the past. This point is discussed in greater detail below. In contrast to the different temperature profiles for QA2, and Q,,,, noted above, those for the reduction processes C2’ and C3’ are similar, in that both pass through maxima at -370 K (- 1OO’C). Furthermore, Qcl, is greater than Qc-*, which is contrary to the behaviour exhibited by the conjugate arrests A2’ and A3’ (Qazs < Qaj, T < 460 K). These relationships provide additional evidence that the potential arrests A2’, C2’ and A3’, C3’ do not arise from simple conjugate charge transfer reactions as implied by most mechanisms that have been proposed to account for the electrochemical behaviour of the silver electrode in alkaline solution.
,
I
571
I
QA/%
A3’/A2’
_
Fig. 8. Charge
ratio
us tcmpcrature
profilm
for various
arrests
obserwd
during
galvanostatic
silver in 1 mol kg-’ KOH. Charging current = 5.7mA cm-*.
charging
of
572
BRUCE G. POUND, DIGBY D.
Ag + 20H2AG + 20H-,
AgzO
+
201~
-
*’Ai’ AgOAZ.AL?
A c2.c2,
-G A3~y
+ H,O f e-
Ag,O + H,O + 2eAg,O,
MACDONALD
(10) (11)
+ Hz0 + 2e- (12)
Thus, if the arrests A2’, C2’, and A3’, C3’ arise from conjugate processes, as described above, then QA3./QA2. and Q&Q,--, should both be equal to unity, assuming complete oxidation of Ag,O to Agz02, in contrast to the complex variations with temperature for both quantities shown in Fig. 8. It can be argued that oxidation of Ag to Ag,O, or even directly to Ag,Oz, can occur in the region of potential arrest A3’, might be expected to be greater than so that QAY/QAZ’ one. This is indeed the case for temperatures less than WK, but it is not so for the highest temperature (478 K). Also, no obvious explanation is apparent for the maximum in this ratio at 370 K. Likewise, Ag,O, could continue to be reduced in the region of potential arrest C2, thereby giving rise to a ratio Qc2’/Qc3. that is greater than one, as observed. Miller[19] found in his rotating ring-disk electrode study under ambient conditions that on reduction of Ag,O, the surface is almost immediately covered with Ag,O. His results indicate a dissolution-precipitation mechanism for this process of the type A&O, 2AgO-
+ 2e- + 2AgO-
(dissolution)
+ H,O * Ag,O + 20H-
(13)
(precipitation)(l4)
Since not all of the dissolved intermediate may pra cipitate out as Ag,O, the net effect is that the charge due to reduction ofAg,O should be less than indicated by that for the reduction ofAg,O, ; that is, Qcz,/Qcs, is expected to be less than unity rather than in excess of one as observed. Accordingly, the dissolutionprecipitation mechanism proposed by Miller[l9] does not appear to account for the variable temperature results report in this study. Finally, the corn$ lexity of the oxidation-reduction mechanism for silver is shown by the variable sweep rate cyclic voltammetry data plotted in Figs. 2 to 5. Thus, the fact that some of the E, us In(u) and i, 0s vl” correlations exhibit nonlinearity, and particularly that the i, us ul” plots in general do not pass through the origin, indicates that the peaks observed by cyclic voltammetry do not arise from simple charge transfer reactions at least over the range of conditions employed in this work. Clearly, a much more detailed study of the kinetics of the electrochemical oxidation and reduction of silver and its oxides is required if the mechanisms for these important processes are to be delineated. SUMMARY AND CONCLUSIONS
The electrochemical behaviour of polycrystalline silver in 1 mol kg-l KOH over the temperature range 295 K to 478 K has been investigated using cyclic voltammetry and cyclic galvanostatic charging. The findings of this study are summarized as foIlows : (1) The processes that are generally assumed to occur
AND
JOHN W. TOMLWSCIN
on the anodic oxidation of silver (Ag + AgO-, Ag -+ Ag,O, Ag,O -, Ag202) are consistent with the calculated thermodynamic equilibrium behaviour of silver in 1 mol kg- ’ KOH over the temperature range of interest. (2) Detailed analysis of the charge associated with each charge transfer process indicates that the processes that give rise to the potential arresta A2, C2 and A3, C3 are not strictly conjugate; that is the reduction does not take place via the simple reversal of the forward oxidation step. The complexity of the redox mechanism is indicated by the nonlinearity of peak potential us In(u) and peak current us ti”’ plots do not pass through the origin. (3) The total anodic charge capacity of polycrystalline silver in 1 mol kg-’ KOH was found to pass through a maximum at - 370 K (- 100°C), thereby indicating that the performance of a polycrystalline silver anode is optimal at this temperature.
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The electrochemistry
of silver in KOH at elevated temperatures
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