A. ET. VIJH
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-36 t
I
SEC1
I
I
Time.
Figure 7. Change in absorbance for supporting electrolytos at applied potentials.
An additional interesting phenomenon is observed when these thin films are used as electrodes in contact with electrolytic solutions. That is, the IRS absorbance changes with applied potential in the absence of any electrolysis of electroactive species. For example, a change, similar to that noted previously for tin oxide filmsj2is shown in Figure 7 for a tin oxide film in contact with an acetate buffer solution when a square-wave potential pulse of $0.6 V vs. reference saturated calomel electrode is applied. The rate of change of absorbance lags considerably behind the rise and fall of the poten-
tial. The maximum absorbance change again occurs at the maximum of the “interference” peaks. Similar changes also occur with glass coated with vapor-deposited gold films, with the only difference being that the absorbance changes are much faster than with tin oxide. It is important that these changes affect the magnitude of the absorbance only and not the wavelength of the “interference” maximum. This suggests that the refractive index is not being changed in the immediate vicinity of the film surface. Perhaps optical rotation of the penetrating electric vector at the film solution interface due to the applied potential can explain these observations. Further work is necessary to understand these results.
Acknowledgment.
The fruitful discussions with
W. N. Hansen of the North American Aviation Science Center are greatly appreciated. The assistance of
R. Chang during the early part of this work is acknowledged. The authors gratefully acknowledge the support of this work by Grant No. G X 14036 from the Research Grant Branch of National Institutes of General Medical Sciences, National Institutes of Health, and by Navy Ordnance Laboratory Contract N123(62738)56006A.
Electrolytic Hydrogen Evolution Reaction on Aluminum in Acidic Solutions by A. K. Vijh R h D Laboratories, Sprague Electric Company, North Adams, Massachusetts 01247 (Received J u l y 31, 1967)
The mechanism of the electrolytic hydrogen evolution reaction (h.e.r.) has been studied on aluminum in several acidic solutions. Experimental measurements consist of galvanostatic current-potential relationships, open-circuit decay from initial cathodic potentials, cathodic-charging curves, and determination of apparent heat of activation. Mechanistic conclusions are based on Tafel slopes, exchange current densities, reaction orders, apparent heat of activation, absence of arrests in the charging curves and potential-decay profiles, nature of capacity-potential curves calculated from open-circuit decay profiles, and some general considerations, e.g., melting point and heat of atomization of aluminum. Significance of other approaches for determination of mechanism of h.e.r. in relation to aluminum is briefly discussed. Initial discharge step is suggested as the likely rate-determining stage (rds) in the over-all reaction; this inference, however, is not entirely conclusive, owing to the difficulties involved in distinguishing initial discharge step from the electrochemical desorption step as the probable rds.
I. Introduction In the present investigations, an attempt has been made to examine the mechanism of the hydrogel1 evolution reaction (h,e.r.) on aluminum in acidic solutions. The solutions in which the h.e.r. has been studied are Oa2 H2S04,Oa5 H2S047o‘s6 HzS04, la7 H2S04J 1 N CH3COOH, and (2 N CH3COOH -I- 1 N NH4COThe Journal of Physical Chemistry
OCH3). The only previous results available in the literature are the galvanostatic evaluation of the Tnfel slopes for the h.e.r. on aluminum in 1 h’ HC1’ and in 2 N H2S04.2 These studies, however, were carried out (1) A. Hickling and F. W. Salt, Trans. Faraday Sac., 36, 1126 (1940).
ELECTROLYTIC HYDROGEN EVOLUTION REACTION ON ALUMINUM before the germination3r4of modern procedures of sohtion purification, electrode preparation, luggin probes, etc. The quantitative significance of these reports,lJ therefore, is rather limited. Aluminum is a typical example of corrodible metals on which not many investigations involving h.e.r. have been carried out.
11. Experimental Section A conventional three-compartment Pyrex cell and the circuits usually involved in the electrochemical measurements were The cell was custommade by the IZontes Glass Co., Vineland, E.J. The three compartments of the cell could be isolated by means of solution-sealed stopcocks. The cell was sealed off from the atmosphere by means of Tru-bore tubing and standard joints. The cell was provided with appropriate gas inlets and outlets. It was cleaned with chromic acid and rinsed several times in deionized, distilled water. Solutions were made up of water which had been deionized and then distilled and CP sulfuric acid or ACS acetic acid, or ACS acetic acid plus ACS ammonium acetate. Purification of the sulfuric acid solutions was carried out by preelectrolysis for 16 hr at the highest current density used in the experiment on a sacrificial platinized-platinum electrode. Preelectrolysis of the acetic acid solutions, however, was carried out only for about 2 hr, since prolonged preelectrolysis seemed to introduce some impurities, as manifested by very irreproducible results, unsteady potentials (at constant current), anomalous Tafel slopes, and increased unsteadiness of potentials with increased rate of bubbling. These impurities, in acetic acid solutions, could arise during preelectrolysis in the anodic compartment and then migrate to the working compartment, since the stopcock between the two compartments had to be slightly open in order to pass the high currents desired during preelectrolysis. Shorter preelectrolysis or even no preelectrolysis at all seemed to give reproducible and steady results in acetic acid solutions. All solutions were also preelectrolyzed on large platinized platinum electrodes (at very low current densities of course) for about 3-5 hr in order to purge the solution of any organic impurities through chemisorption on the large electrode surface. Prepurified hydrogen was bubbled through the working compartment throughout preelectrolyses. No special claims are made regarding the purity of solutions, except that the residual steady-state contamination, if any, seemed not to interfere with the reproducibility and steadiness of the results; also, there was no effect of the rate of bubbling on the values of potential (at constant current) observed. This purification procedure when coupled with proper electrode surface preparation (to be described below) seemed to give reproducible Tafel plots over several decades of current density and without anomalously high Tafel slopes.
1149
All experiments except those for the determination of the apparent heat of activation were carried out at room temperature (25 1 1.5’). The apparent heat of activation was determined in 1 N CH3COOH solutions from the exchange current densities obtained at 1, 12.5, 30.5, 38.5, and 55.5’. The entire cell was suspended in a large water bath with all the three compartments and the solution-sealed stopcocks dipping in the bath well above the level of the solution within the cell. S o elaborate apparatus for temperature control of the bath was used. The temperature was controlled manually and the temperature stayed well within 10.25’ of the stated values during measurement of a given Tafel relation. Before the commencement of the runs, the solution was thoroughly purged of oxygen and saturated with prepurified hydrogen. The working electrodes were prepared by engaging an aluminum wire (99.98%) into a glass tube by means of a joint made up of a special Teflon tubing, Flotite, manufactured by Pope Scientific, Inc. Such a joint is mechanically sound, leakproof and chemically inert toward strongly acidic polishing baths and degreasing organic solvents. A mechanical seal between glass-Teflon-aluminum’ was also found to be satisfactory, except for the awkward electrode size in some runs in which very high current densities were desired for a given current. All working electrodes mere chemipolished for 2 min at 90” in a “bright-dip” bath (85% H3P04 and 15% HS03) and were subsequently mashed with deionized, distilled water, dipped in 1 N NaOH (10 min) and then washed again several times in deionized, distilled water. This procedure was carried out just before the commencement of the run and seems to give reproducible Tafel slopes over several decades of current densities. The counterelectrodes were made up of smooth platinum wire. Hydrogen reference electrodes were used in all experiments. During all experiments, the three compartments were isolated by wet but closed solution-sealed stopcocks. 111. Results A. Steady-State Current-Potential Relationships. In Figure 1,results are presented for potential-log (current density) relationships for h.e.r. on aluminum in 0.2 N HzS04 (PH 0.94), 0.5 N HzS04 (PH 0.59), 0.86 N HzS04 (pH 0.37), and 1.7 H2S04 (pH 0.1); the results are compared with some previous ones2in 2 N HzS04(pH 0.04), Tafel slope, b, and exchange current density, io,have the values equal to 0.11 V and A cm-2, respec( 2 ) A. G. Pecherskaya and V. V. Stender, Zh. Prikl. Khim., 19,1302 (1946). (3) J. O’M. Bockris, Chem. Rev., 43, 525 (1948). (4) J. O’M. Bockris, Trans. Faraday Soc., 43, 417 (1947). (5) B. E. Conway and M. Dzieciuch, Can. J . Chem., 41, 21 (1963). (6) J. J. hIacDonald and B. E. Conway, Proc. Roy. Soc. (London), A269, 419 (1962).
(7) N. D. Greene, “Experimental Electrode Kinetics,” Rensselaer Polytechnic Institute, Troy, N. Y., 1965
Volume 72, Number 4 A p r i l 1968
1150
A. K. VIJH
-.7
-
..6
-
->
I
-6
I -.5-
-.4
-
-.3
-
-21
16‘
16’
16‘ RATE
,
16’
idz
I
16~
.~ ~
I
16’
(a.cr?)
16‘
16‘
16~ RATE
Figure 1. Tafel plots for h.e.r. on A1 in HzS04: 0, 0.2 N; 0, 1.7 N; and a, results of Pecherskaya and Stenderz in 2 N HeS04. Values of b and log ioare 0.11 and -7.7, respectively.
0, 0.5 N ; A, 0.86 N;
tively. It may be pointed out that the rate at a given potential seems to be independent of pH in the p H range 0.1-0.94. However, the scatter in this graph is almost comparable to the changes in rate, a t a given overpotential, expected in going from pH 0.1 to 0.94. Increased scatter at lower cathodic potentials probably arises from complications due to “chemical” corrosion reactions or from the presence of trace impurities in the solution. Some data of Pecherskaya and Stender,2 also shown in Figure 1, seem to depart noticeably from the Tafel line drawn through our results for the four concentrations of sulfuric acid studied (0.2, 0.5, 0.86, and 1.7 N). This probably arises from the questionable purity of the solutions of these authorsj2which has also been commented upon previo~sly.~ Tafel plots in 0.86 N H2S04 (pH 0.37) have been compared with those in 1 N CH3COOH (pH 2.4), and in 2 N CH3COOH 1 N iSH4COOCH3(pH 4.4) in Figure 2. These are the results from which various reaction-order derivatives10-12 have been determined. The derivation of reaction orders from these plots is justified because (a) these are all concentrated solut i o n with ~ ~ ~a consequently constant value of potentials;lOvll (b) the reaction orders have been determined only from potentials equal t o or more cathodic than -0.4 V, which is roughly the potential of zero charge for aluminum.14 At these high cathodic potentials, any complications that otherwise might arise from the specific adsorption of acetate ions are vanishingly small. Absence of specific adsorption of acetate ions is also suggested by the fact that there is no increase in the value of Tafel slopeloin acetate solutions, even at lower temperatures (see Table I). Hence for reaction-order purposes, the Tafel plots in different sulfuric acid and acetic acid solutions may simply be assumed, t o a first approximation, as potential-rate relations obtained in solutions of different pH values (ie., 0.37,2.4,and4.4). This direct comparison of potential-rate relations in solutions of different chemical composition is further
+
+
The Journal of Physical Chemistry
t
16~
162
16’
(a.ci?l
Figure 2. Tafel plots for h.e.r. on A1 in acidic solutions of different pH values: 0, pH 4.4, b = 0.122 V (2 N CHaCOOH f 1 N NH~COOCHI);0, pH 2.4, b = 0.117 V (1 N CHICOOH); A, pH 0.37, b = 0.114 V (0.86 N HzS04).
Table I: Aluminum in 1 N CHgCOOH (h.e.r,)
Temp,
Temp,
b,
1000(1/T),
OC
OK
V
io, Acm-2
Rest potential,
TinOK
V
55.5 38.5 30.5 12.5 1.0
328.5 311.5 303.5 255.5 274.0
0.124 0.118 0.117 0.123 0.125
-8.15 -8.60 -8.82 -8.86 -9.28
3.022 3.214 3.296 3.505 3.650
0.430 0.388 0.368 0.312 0.265
justified by the following facts. (a) The species involved in the discharge step in all three solutions must be the same, i.e. H 3 0 + and H30+ alone. (b) The Tafel slope in all the three solutions is the same, i.e., 2.3(2RT/F) (118 k 4 mV). (c) Identical transient behavior is observed in the three solutions in the charging curves and in the open-circuit decay profiles. It may seem rather arbitrary that only one sulfuric acid solution (0.88 N ) has been chosen for the reactionorder purposes (Figure 2 ) . However, this is the only one which is fairly concentrated with a consequently constant potential10,l1and at the same time does not
+
(8) M. Pourbaix, “Atlas D’Equilibres Electrochimiques,” GauthierVillars and Co., Paris, 1963, p 168. (9) J. O’M. Bockris in “Electrochemical Constants,” National Bureau of Standards Circular 524, U. S. Government Printing Office, Washingtan, D. C.,1953. (10) B. E. Conway, “Theory and Principles of Electrode Processes,” The Ronald Press Co., New York, N. Y., 1965. (11) B. E. Conway and M. Salomon, Electrochim. Acta, 9, 1599 (1984). (12) K. J. Vetter, “Electrochemische Kinetik,” Springer-Verlag, Berlin, 1961; see also E. Yager, Ed., “Transactions of the Symposium on Electrode Processes,” John Wiley and Sons, Inc., New York, N. Y., 1961, p 47. (13) A. N. Frumkin in “Advances in Electrochemistry and Electrochemical Engineering,” Vol. I, P. Delahay, Ed., Interscience Publishers, Inc., New York, N. Y., 1961, p 74. (14) L. I. Antropov, “Kinetics of Electrode Processes and Null Points of Metals,” Council of Scientific and Industrial Research, New Delhi, India, 1960.
1151
ELECTROLYTIC HYDROGEN EVOLUTION REACTION ON ALUMINUM -.4v
-.IV
-.6V
-.7v
lo-'
1
1
I
0
2
I
1
5
3 p H 4
6
,$, Figure 3. Plot for reaction order, ( b In i / b In C H ~ O + ) ~for aluminum in solutions of different pH. Data for the four values of TJ shown in this figure have been obtained from Figure 2; (a In i / b In C H ~ O + ) ~=, $0.39.
1
I
I
I
Figure 5. Plot for ( b In i / b In C H ~+)+.+ O for aluminum in acidic solutions: data calculated from Figure 2 (see text); , 4 = -0.65 V. Data obtained from 0, 4 = -0.5 V; . Figure 2; value of (a In i / b In CH~O+)+,$ obtained is 0.87.
I
cause as drastic a chemical corrosion reaction as 1.7 N sulfuric acid. I n any case, a visual examination of the reaction-order plots (Figures 3-5) shows that the values of various derivatives would not change significantly if the Tafel plots obtained in any of the other sulfuric acid solutions (ie., Figure 1) are used in Figures 2-5, instead of the one obtained in 0.86 N sulfuric acid. It may be mentioned that anomalously high Tafel slopes (>0.13 V ) were observed when contamination of the solutions and/or electrode surface was suspected by facts such as vacillating potentials (at constant current),
irreproducible rates, very short linear Tafel regions (about one decade), strong susceptibility of the potentials to rate of bubbling in the working compartment and, finally, insensitivity of the electrode potential to A cm-2) low polarizing current densities (ca. 1 X thereby suggesting impurity reactions. These results have been discarded. Repeatedly, linear Tafel regions over about three decades with a Tafel slope close t o 2.3 (2RTIF) were obtained, to the exclusion of aforementioned symptoms of contamination, on freshly prepared electrode surfaces in solutions purged of the obvious impurities, e.g., dissolved oxygen. The results were particularly susceptible to impurities diffusing from the anodic compartment, as expected; this situation was remedied by keeping the solution-sealed stopcock between the anode and the cathode compartments always closed during the measurements. B. Reaction Orders. In Figure 3, plots of pH vs. rate of h.e.r. a t the four indicated values of electrode potential measured against reversible hydrogen electrode in the same solution are shown and have been derived from the data presented in Figure 2. The value of the derivative (b In i / b In C W ~ O +as) obtained ~ , ~ ~ ~ ~ ~ ~ from these pH-rate relationships is 0.39 for every one of the four electrode potentials shown in Figure 3. I n Figure 4, relationships between pH and overpotential (same as the electrode potential in case of h.e.r.), q , measured against reversible hydrogen electrode in the same solution a t two shown values of constant current density have been presented. These relations, again, are derived from the data in Figure 2. Volume 76, Number 4 April 1968
A. K. VIJH
1152
The value of the derivative (a& log C H , O + ) ~ , + , ~ ~ . ~ ~ estimated from these plots, for either of the current densities shown, is equal to 45 mV. An attempt has been made to determine chemically significant reaction orders in Figure 5 by plotting us. rate for the h.c.r.; 4 is the value of the electrode potential against standard reversible hydrogen electrode. The data used in Figure 5 have been calculated from results in Figure 2 and from a known dependence of potential of standard reversible hydrogen electrode on pH, i.e., 59 mV/pH unit'QJ1at room temperature. The value of this chemically significant reaction order, (a In i/d In CH,~.)+.+, is 0.57. C. Nonsleudy-Slafe (Transient) Sludies. A typical transient (in 0.2 N HzSOI) depicting the electrode potential-time behavior obtained on cessation of initial cathodic p0larization~~~~'~~'2.'5.1" has been shown in Figure 6. Absence of any arrests or inflections, which usually are diagnostic of open-circuit desorption of a JI i l . -.u) a. 4 8 E" I V ) species (the possible species in this case is adsorbed H), Figure 7. A capacitance-potential profile calculated from deposited in a previous steady-state polarisaexperimental data represented in Figures 1 and 6; tionS.6J0J2.'SJ8may be noted. In Figure 7, capacity"rest" potential in this case is ca. -320 mV. electrode potential relationship has been presented; this graph is based on the open-circuit potential-decay profiles shown in Figure 6 and steady-state currentdecay profiles are closer to the "equilibrium" values potential relationship for 0.2 N H2SO4shown in Figure and hence more significant than the e values that may 1. In Figure 7, absence of a characteristic capacity be calculated from the charging curves or from profiles maximum (sharp for the Langmuir case and flat for the depicting the forced discharge behavior of initial steadyTemkin case)*OJ7may be noted. Another point to be state potential.l* emphasized here is the value of the maximum capacity, A typical cathodic-charging curve (in 0.2 N &SO,) which is remarkably lower than that estimated for has been shown in Figure 8. Any arrests or inflection coverage by adsorbed H approaching unity." The in the potential-time curve which can be diagnostic of magnitude of e observed in Figure 7, however, is greater the adsorbed species,ls are again absent. Absence of than the usual values for double layer capacity and inflections in the charging curves has been observed for probably indicates some pseudo-capacity associated all values of the cathodic-charging current, uiz. from with potential dependence of charge transfer involved 1 X 10-~to 2 X lo-' A Identical results have in the chemical corrosinn reaction at low cathodic pobeen obtained in acetic acid solutions. tentials. It may be mcntioncd that thevalues of cepacD . Apparent Heat of Acliualion. Tafel relations ity, c, associated with potential dependence of any have been obtained galvanostatically on aluminum in possible coverage ohtaincd from open-circuit potential1N CH,COOH solutions a t 1, 12.5,30.5,38.5,and 55.5" and are shown in Figure 9. These relations were obtained from descending temperatures after a steady "corrosion potential" had been achieved at 55.5'. The Tafel slopes, exchange current densities, and "rest" potentials (Le. the mixed corrosion potentials) have been shown along with the corresponding values of tempera.O. 92 V ture, T (in "I
