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Hydrogen evolution on oxide electrodes: Co3O4 in alkaline solution
255
.I. Electroanal. Chem., 339 (1992) 255-268
Elsevier Sequoia !&A.,Lausanne
JEC 02074
Hydrogen evolution on oxide electrodes: Co,O, in alkaline solution * E. Veggetti,
I.M. Kodintsev
l
* and S. Trasatti
l
**
Department of Physical Chemistry and Electrochemistry, University of Milan, via Venezian 21, 20133 Milan (Italy)
(Received 11 February 1992)
Abstract
Mechanistic and electrocatalytic aspects of H, evolution from alkaline solutions were investigated using Co,O, cathodes prepared by thermal decomposition of cobalt nitrate. Two different arrangements were used: Ti/Co,O, (A) and Ti/RuO, /C&O, (B). The surface response was investigated by cyclic voltammetry, while the kinetics was studied by determining Tafel lines and reaction orders. The results show that electrodes B are more active than electrodes A with an overpotential decrease of more than 0.2 V. The mechanisms are also different: the reaction order is fractional with electrodes B. The experimental picture for B electrodes is consistent with H, evolution taking place on RuO, at the bottom of pores and/or cracks in the Co,O, layer. Presumably, the latter oxide is cathodically reduced to some form of hydroxide which is not particularly active for H, evolution. In addition, the oxide layer is subjected to mechanical erosion due to the action of the evolving gas.
INTRODUCTION
Oxide electrodes order
to improve
(DSA@) were the performance
introduced into electrochemical technology in electrodes [1,2]. of gas CO,, Cl,) evolution
Ernest Yeager was among the first to investigate their electrocatalytic properties, particularly for 0, evolution and reduction [3]. Recent patents have also claimed good performances for DSA-type electrodes (including RuO,, 110, and spinels) in
Dedicated to Professor Ernest B. Yeager on the occasion of his retirement and in recognition of his contribution to electrochemistry. l * On leave from Mendeleev Institute of Chemical Technology, Moscow, Russian Federation. l ** To whom correspondence should be addressed.
l
0022-0728/92/$05.00
0 1992 - Elsevier Sequoia S.A. All rights reserved
256
the cathodic evolution of hydrogen [4-71. However, there has been little fundamental research in this field. Only a few papers in the open literature deal with the study of hydrogen evolution at oxide electrodes [8-101. The first investigation was carried out in this laboratory long ago [ill before any practical implication was envisaged. Oxides are thermodynamically unstable in the hydrogen adsorption region [12,13] and are thus in principle not expected to be good electrocatalytic cathodes. A prerequisite for understanding the good performances obtained in practice is to elucidate the kinetic mechanism and the surface response. In previous work, we have investigated the electrocatalytic behaviour of IrO, and RuO, cathodes in acid solutions [14-161. In this paper, results of a study of C&O, in alkaline solutions are reported.
EXPERIMENTAL
Electrodes C&O, layers on a Ti support were prepared by thermal decomposition of Co(NO,), .6H,O (Fluka) at temperatures in the range 230-500°C. The precursor was dissolved in isopropanol so as to give a 0.3 mol dmW3 solution and spread onto the support by brushing. After each operation, the solvent was evaporated at 60-70°C and the dried layer was fired for 10 min in a oven preheated to the required temperature. The operation was repeated until the desired oxide loading (as determined by weighing) was reached. Electrodes were finally calcined at the same temperature for 1 h. The typical electrochemical surface properties of Co,O, electrodes prepared in this way have been described elsewhere [17]. The average oxide loading was 3.6 f 0.2 mg Co,O,, corresponding to a nominal thickness of about 3 pm. The mounting of the electrodes and the special Teflon holder have been described elsewhere [18]. Co,O, was prepared at 230, 260, 300, 350, 400, 450 and 500°C. Two samples were prepared at each temperature. Support Ti platelets of dimensions 10 mm X 10 mm X 0.2 mm were used as a support. The metal was first sandblasted and then etched with boiling 10% oxalic acid for 10 min (just prior to depositing the oxide layer) in order to remove any insulating TiO, film. In order to minimize the ohmic drops related to TiO,, which interfere with kinetic determinations, a set of samples fired between 400 and 500°C was prepared by placing an interlayer of RuO, between Ti and Co,O,. A 0.02 mol dmV3 RuCl, * nH,O solution in isopropanol was used to prepare a thin RuO, layer by thermal decomposition at the same temperature.
Solution
Aqueous solutions of 1 mol dme3 NaOH prepared volumetrically were used throughout. They were deaerated before each run and stirred by bubbling nitrogen. The three-compartment cell has been described elsewhere [19]. Experiments were conducted at 25 f O.l”C in a water thermostat. Reference electrode
Potentials were measured and are reported on the saturated calomel electrode (SCE) scale. The reference electrode was connected to the cell through a Luggin capillary to minimize uncompensated ohmic drops. Techniques The surface response was investigated by means of voltammetric curves recorded at 20 mV s- ‘. The kinetics of hydrogen evolution was studied by using quasi-stationary current-potential curves. Currents were read after 3 min at each potential. Potentials were increased in 10 mV steps up to a current of approx. 0.1 A cm-’ and then reversed. Any hysteresis vanished after the second cycle. The order of reaction with respect to OH- was determined by varying the NaOH concentration between 0.1 and 1 mol dme3 while keeping the ionic strength constant with additions of NaClO,. AMEL instrumentation was used throughout. RESULTS
AND DISCUSSION
Voltammetric curves
Voltammetric curves were used to monitor the surface state of the electrodes and to verify the presence of uncompensated ohmic drops related to TiO, films.
Fig. 1. Typical voltammetric curve obtained between -0.1 NaOH for a Ti/Co,O, electrode prepared at 260°C.
and 0.5 V SCE at 20 mV s-’
in 1 mol drne3
258
KXE)
Fig. 2. Voltammetric curve obtained between -0.1 and 0.5 V SCE at 20 mV s-’ in 1 mol dm-3 NaOH for Co30, electrodes prepared at 500°C: (1) Ti/Co,O,; (2) Ti/RuOz /Co,O,.
“Standard” curves were recorded between -0.1 and 0.5 V (just prior to oxygen evolution). Figure 1 shows a typical voltammetric curve for Co,O, in alkali. Details have been given elsewhere [ 171. The curve is characterized by a pair of peaks around 0.4 V related to the Co(III)/cO(IV) redox transition. This surface transition is reversible and therefore the anodic and cathodic peaks should coincide. The separation between the two peaks can be taken as proportional to the uncompensated ohmic drop. Up to 350°C the separation AE between the peaks is small and this indicates the absence of large ohmic effects. At higher temperatures the curves are distorted. However, as Fig. 2 shows, a typical curve is again observed at these temperatures if some RuO, is present in the interlayer. Therefore the distortion is attributed to the insulating properties of TiO,, which is heavily doped by RuO, when an interlayer is present. Voltammetric charge The voltammetric charge q * obtained by integration of the voltammetric curves is determined by the (formal) redox transition [17] CoO(OH)
+ Coo, + H++ e-
(1)
The charge measures the number of surface active sites and can therefore be taken as proportional to the true surface area.
259
8 I
300
I
I
. . 1
400
. I
500
t/C
Fig. 3. Variation of the initial voltammetric charge 4: with the calcination temperature: A Ti/RuO, /Co,O,.
0 Ti/Co304;
Figure 3 shows the variation of q: with the calcination temperature. q* is seen to decrease monotonically as expected for a surface area effect. At T 2 400°C q* drops for the samples without RuO,; however, this is interpreted as an artefact resulting from the distortion of the voltammetric curves. In fact, the qi* values for the RuO,-containing samples at 400°C match the set of values obtained for the other electrodes. Thus a continuous variation of qi* is obtained over the temperature range investigated. Effect of the negative potential limit If the negative limit of the potential range is progressively lowered, the voltam-’ metric curve is dramatically modified. As the potential enters the hydrogen evolution region, a prominent anodic peak develops around - 0.7 V. This is shown in Fig. 4. Comparison with results in the literature obtained with Co(OH), electrodes [20-221 suggests that the onset of atomic hydrogen formation catalyses the reduction of the oxide layer with penetration of hydrogen into the bulk of the oxide. This is an important difference with respect to RuO, and IrO, electrodes which do not exhibit major modifications of the voltammetric curve upon hydrogen evolution 114,161.It is very likely that thermal Co,O, is hydrogenated to hydroxide. The modification is irreversible because the voltammetric charge starts increasing after the imposition of a cathodic load. This aspect will be discussed in detail later on.
260
Fig. 4. Effect of hydrogen evolution on the voltammetric curve at 20 mV s-’ in 1 mol dmm3 NaOH. The electrode was held at - 1.3 V SCE for (1) 10 s, (2) 30 s, (3) 1 min and (4) 5 min.
It is interesting that the electrodes containing RuO, show a much lower tendency towards modification of the voltammetric pattern. It is probable that the presence of RuO, influences also the behaviour of Co,O, as well as that of the support-oxide interface. Current-potential curves Figure 5 shows E-log j curves in the region of hydrogen evolution for two electrodes, one with (electrode B) and the other one without (electrode A) the RuO, interlayer. The difference is dramatic. While hydrogen evolution starts around - 1.05 V SCE for electrodes B, it does not begin until - 1.25 V for electrodes A and is preceded by a slow reaction (which is assumed to be hydrogen penetration into the oxide) with an apparent Tafel slope of about 180 mV. In both cases the main Tafel line is distorted by uncompensated ohmic drop. Analysis of the deviations [23] from the expected straight line reveals that the actual Tafel slope for electrodes without RuO, (A electrodes) is 60 mV if an uncompensated ohmic component of OS-O.6 51 is assumed. This value is typical of ohmic drops in solution at the Luggin capillary. The fact that the same value is obtained at all temperatures suggests that any TiO, layer is presumably reduced on cathodic load and no longer presents a significant ohmic resistance. It is also interesting to note that the calcination temperature does not appear to have any appreciable influence on either the shape or the position of the polarization curve. In other words, surface area effects do not show up at this stage. The Tafel lines of electrodes with RuO, (B electrodes) are less distorted because they start at lower current densities. The Tafel slope at low overpotentials is very close to 40 mV, and is the same for different calcination temperatures. No surface area effect can be detected in this case. This suggests that the active
261
1
I
-4
I
-3
I
I
-1
lo&/A cnY2j2
Fig. 5. Quasi-stationary current-potential curves for H, evolution on Co,O, electrodes: (1) Ti/Co,O,; (2) Ti/RuO* /Co,O,. Tafel lines from ohmic drop corrections are shown. The slopes of the lines are indicated.
surface taking part in the reaction does not depend on the calcination temperature. In the high overpotential region of B electrodes, the analysis of ohmic drop deviations reveals that a second Tafel line with a slope close to 120 mV is present. The ohmic component turns out to be slightly higher with this type of electrode, being about 1.0 fl. Reaction order The current was measured at - 1.24 V for electrodes without RuO, (A) and at - 1.07 V for electrodes with RuO, (B) so as to stay in the linear section of the Tafel line as the pH was changed. Figure 6 clearly shows that the reaction order is definitely different, in the two cases, with a value of - 1 (the reaction rate is depressed by increasing NaOH concentration) for electrodes A and a value of - 0.5 (fractional) for electrodes B. Fractional reaction orders have previously been observed for H, evolution with IrO, and RuO, in acid solutions [14,1.5]. H,-induced modifications H, evolution induces modifications on Co,O, electrodes irrespective of the presence or the absence of .the RuO, interlayer. Immediately after the runs shown in Fig. 4, the “standard” voltammetric curves revealed an increased value of charge. Figure 7 shows that this charge is proportional to the initial charge
262
E >
(4
\ x0
m-
j
\
\
\
\
l’ \
a2t 13
13.4
\
'e
\
\ 13.8
Fig. 6. Dependence of the H, evolution reaction rate at constant potential on the solution pH: (a) Ti/Co,O,, E = - 1.24 V SCE, (b) Ti/RuO, /C&O,, E = - 1.07 V SCE. Straight lines (- - - - - -) of slope (a) 1 and (b) 0.5 are shown.
obtained with fresh electrodes [16]. Thus only subsurface sites are activated by a moderate H, discharge so that a “memory” of the surface area ratio between different samples is still retained. After extensive H, evolution (i.e. after determination of the Tafel line and reaction order) the charge becomes approximately independent of calcination temperature. This is shown in Fig. 8 where the charge after recording the curves in Fig. 4 is plotted against the final charge. Although some scatter is observed, the final value of the charge is about 160 mC crne2. It should be noted that this value is more than five times higher than the highest initial value. Since 1.8 mg cmW2 corresponds to 7.5 pmol cmA2 Co,O,, th e charge associated with (roughly) one electrode exchange per Co atom amounts to about 2 C cmm2. This is approximately 10 times higher than the final charge effectively measured. This may indicate that complete bulk reduction does not take place, and that only the total surface (“outer” + “inner”, i.e. including cracks, pores and grain boundaries), with an approximate ratio of 1: 10 between surface and bulk atoms, is reduced. However, two facts have to be borne in mind.
263
/ 40-
/ l
k
/
/
/
::’
A
20-
/
/
’ k
,$ ,d
i I
10
I
I
20 30 4r/mC cm-*
I
40
Fig. 7. Plot of the voltammetric charge 4: (between -0.1 and 0.5 V SCE) from Fig. 2 for freshly prepared electrodes vs. the voltammetric charge q: after recording the curves in Fig. 4: - - - - - straight line with unit slope; l Ti/Co,O,; A Ti/RuO, /Co30,.
(1) The voltammetric curve is recorded at a finite sweep rate (20 mV s-l). If hydrogen penetration is slow, the diffusion front can penetrate only marginally during a voltammetric cycle [24]. (2) Erosion of the overlayer is visible to the naked eye. The measured final charge could refer to a much smaller thickness while involving reduction of the whole layer. The latter explanation seems the more probable on the basis of the complete experimental picture. If the final charge still involved only surface atoms, it would be difficult to understand how it can be independent of the calcination temperature. Since the particle size differs at different temperatures, a difference in surface area must persist even after extensive surface wetting promoted by H, evolution. This particular aspect deserves further experimental investigation. However, it is also possible that complete “hydrogenation” of the C&O, is a very slow process, and
264
“i .
JpJo- ‘.I
.
z ... .: l
\‘6
. .
.
r.., 50
100 c@nC cn?
150
Fig. 8. Plot of q: (from Fig. 7, same symbols) vs. the voltammetric charge after extensive H, evolution (q%)(at the end of all experiments).
what is seen in this work is only its beginning; nevertheless, it is already sufficient to smooth down surface area effects. It is intriguing that a higher value of the final charge is on average obtained for electrodes with an RuO, interlayer. This is probably due to the fact that the measured charge includes that exchanged by the surface of RuO, itself. Electrocatalytic activity It has already been mentioned that the rate of hydrogen evolution does not appear to depend on the calcination temperature. This is clearly shown in Fig. 9 where the apparent current density is plotted versus the final charge in a log-log
0.5 9 : E > j
O-
-0.5l 1.5,
I
I
2.5 log (s’/mC cm? Fig. 9. Log-log plot of the reaction rate for H, evolution against the final voltammetric charge: - - - - - iine with unit slope. Symbols as in Fig. 7. 2D
265
a4-
e,3-
E
A \
> !? 2
\*
0.2-
\ A \ A
0.1 -
\A
\
\
0
-‘-$--,--g--.
1
0
_a---____ *
I
I
I
300
400
!500
t/‘C Fig. 10. Dependence of the reaction rate for H, evolution normalized to unit surface charge on the calcination temperature: (1) Ti/Co,O.,; (2) Ti/RuO, /C&O,.
diagram. Undoubtedly, there is no correlation between j and qf*; the former remains on average constant as the latter changes. More specifically, Fig. 10 shows a plot of the current density per unit surface charge as a function of the calcination temperature. The electrocatalytic activity per active surface site is seen not to vary in any systematic way with the calcination temperature. Kinetic mechanism The results of the kinetic studies are summarized in Table 1. The reaction mechanism is definitely different for the two sets of electrodes. The 60 mV Tafel slope indicates that the rate-determining step is a chemical reaction following the primary discharge. However, the 40 mV Tafel slope suggests that the second
TABLE 1 Experimental kinetic parameters for H, evolution on Co,O, in alkaline solutions Electrode type
Tafel slope b /mV
Reaction order voH-
Ti/CosO., (A) Ti/RuO, /Co,O., (B)
approx. 60 mV approx. 40 mV
approx. - 1 approx. - 0.5
266
electron transfer is rate determining. In the absence of more detailed molecular information, a common mechanism with two different rate-determining steps can be envisaged:
I
-M-OH
+ H,O + e-
I
-M-OH,*
I
+ e-
I
-M-H+H,O
-M-OH:
+ OH-(aq)
(2a)
I
-
-M-OH,
I
-
-M-OH,
P) I
-
-M-H
I
-
-M-O
PC)
+ OH-(aq)
I
+ H, t
(2d)
The asterisk indicates an unstable surface species which needs to rearrange to a more stable species to become active for H, evolution. Results indicate that H, evolution on Ti/Co,O, electrodes (A) is preceded by an anomalous Tafel region which might describe the slow primary reduction of surface sites (b = 120 mV> with a transfer coefficient lower than 0.5 (based on the assumption of a dual barrier model). H, evolution on these electrodes is probably limited by the slow rearrangement of the reduced surface sites, i.e. step (2b) (b = 60 mV), so that steps (2~) and (2d) only take place at much higher overpotentials. With Ti/RuOJCo,O, electrodes both step (2a) and (2b) are fast so that step (2~) becomes rate determining (b = 40 mV). The proposed mechanism is also consistent with the observed reaction order. Previous results [25] have shown that the surface charge of oxides is pH dependent so that the electric potential at the reaction site varies by 59 mV per pH unit: $J* a - (RT/F)
ln[OH-]
If step (2b) is rate determining,
(3) the kinetic equation is
ja[-OH;]
(4)
Since step (2a) is in quasi-equilibrium, [-OH:]
a [OH-]-’
exp( -FE/Z?T)
(5)
From eqns. (4) and (5) j a [OH-]-’
exp( -FE/RT)
which reproduces
the kinetic parameters
(6) in Table 1 for Ti/Co,O,
electrodes.
267
If step (24 is rate determining,
the kinetic equation is
ja[-OH,]exp[-cy,(E+4*)F/RT]
l
(7)
Since [-OH,]
a [OH-]-’
exp( -FE/RT)
(8)
it follows from eqns. (31, (7) and (8) that j
a
[OH-]-“’
exp[ - (1 + cz,)FE/Z?T]
(9) where ac is the cathodic transfer coefficient. Assuming (Y,= 0.5, eqn. (9) reproduces the kinetic parameters in Table 1 for Ti/RuO,/Co,O, electrodes. The Tafel slope close to 120 mV at higher overpotential for B electrodes is probably related to a shift of the rate-determining step from step (2~) to step (2a) (primary discharge). This is quite often observed in hydrogen electrocatalysis mechanisms [ 10,261. In view of the similarity of the voltammetric curves for electrodes with and without an RuO, interlayer, which indicates that the chemical nature of the surface sites is unchanged, the only possible explanation for the different kinetic mechanisms is that H, evolution takes place on the RuO, sites at the bottom of pores, cracks and crevices. A similar situation has been observed with RuO, layers covered with chemically deposited Ni [lo]. This would account for the additional ohmic loss along the pores which is not observed for electrodes without RuO,. CONCLUSIONS The present results show that Ti/Co,O, electrodes are not very active for the cathodic H, evolution reaction in alkaline solution. Under a cathodic load Co,O, is reduced and H, probably penetrates into the oxide bulk. H, evolution takes place on the reduced oxide with a higher overpotential (> 0.2 V) than expected. The reaction rate is limited by a chemical step following the primary discharge, probably involving some rearrangement of the surface intermediate (a reduced surface species). Ti/RuO,/Co,O, electrodes are more active. H, evolution takes place in the expected potential range so that the overpotential gain is > 0.2 V with respect to pure C&O,. The active surface is not that of Co,O, but that of RuO, at the bottom of pores, cracks, loose grain boundaries etc. The kinetic mechanism is different, as the reaction rate is limited by the transfer of the second electron (further reduction of surface intermediates formed by reduction of water molecules). The observed fractional reaction order indicates that a surface charging effect, related to the acid-base properties of the oxide, is operative.
l In the case of metals, double-layer corrections are written as E - c$* since q~* decreases as the ionic concentration increases. In the case of oxides the correction must be written as E + 4* since the pH governs the surface charge and S#J*increases as the concentration of the potential-determining ion increases.
268
The intense evolution of H, gas provokes erosion of the active layer. For this reason the final charge associated with the redox transition of Co,O, is far less than that calculated for the initial catalyst loading. It thus appears that pure Co,O, does not show particularly interesting electrocatalytic properties for H, evolution in alkaline solution, although it may be useful in specific cases. ACKNOWLEDGEMENT
Financial support of this work by the CNR (National Research Council, Rome) is gratefully acknowledged. REFERENCES 1 S. Trasatti (Ed.), Electrodes of Conductive Metallic Oxides, Parts A and B, Elsevier, Amsterdam, 1980, 1981. 2 S. Trasatti and W.E. O’Grady, Adv. Electrochem. Electrochem. Eng., 13 (1980) 177. 3 W.E. O’Grady,‘C. Iwakura, J. Huang and E. Yeager, in M.W. Breiter (Ed.), Electrocatalysis, Electrochemical Society, Pennington, NJ, 1977, p. 286. 4 E. Nicolas, Eur. Pat. Appl. 23368, 1981; Chem. Abstr., 94 (1981) 199838. 5 J. Clerc-Renaud, F. Leroux and D. Ravier, Eur. Pat. Appl., EP 240413, 1987; Chem. Abstr., 108 (1988) 45864. 6 J.F. Cairns, D.A. Denton and P.A. Izard, Eur. Pat. Appl., EP 129374, 1984; Chem. Abstr., 102 (1985) 102442. 7 A. Nidola, PCT Int. Appl., WO 86 03790 (19861; Chem. Abstr., 105 (1986) 122974. 8 E.R. Kijtz and S. Stucki, J. Appl. Electrochem., 17 (1987) 1190. 9 M. Jaccaud, F. Leroux and J.C. Millet, Mater. Chem. Phys., 22 (1989) 105. 10 A. Anani, Z. Mao, S. Srinivasan and A.J. Appleby, J. Appl. Electrochem., 21 (1991) 683. 11 D. Galizzioli, F. Tantardini and S. Trasatti, J. Appl. Electrochem., 5 (1975) 203. 12 M. Pourbaix, Atlas d’Equilibres Electrochimiques a 25°C Gauthiers-Villars, Paris, 1963. 13 S. Trasatti and G. Lodi in S. Trasatti (Ed.), Electrodes of Conductive Metallic Oxides, Part A, Elsevier, Amsterdam, 1980, p. 301. 14 J.F.C. Boodts and S. Trasatti, J. Appl. Electrochem., 19 (1989) 255. 15 J.F.C. Boodts, G. Fregonara and S. Trasatti in F. Hine, B.V. Tilak, J.M. Fenton and J.D. Lisius (Eds.), Performance of Electrodes for Industrial Electrochemical Processes, Electrochemical Society, Pennington, NJ, 1989, p. 135. 16 S. Ardizzone, G. Fregonara and S. Trasatti, J. Electroanal. Chem., 266 (1989) 191. 17 R. Boggio, A. Carugati and S. Trasatti, J. Appl. Electrochem., 17 (1987) 828. 18 R. Garavaglia, C.M. Mari and S. Trasatti, Surf. Technol., 23 (1984) 41. 19 G. Lodi, E. Sivieri, A. De Battisti and S. Trasatti, J. Appl. Electrochem., 8 (1978) 135. 20 L.D. Burke, M.E. Lyons and O.J. Murphy, J. Electroanal. Chem., 132 (1982) 247. 21 H. Gomes Meier, J.R. Vilche and A.R. Arvia, J. Electroanal. Chem., 138 (1982) 367. 22 L.D. Burke and M.M. Murphy, J. Electrochem. Sot., 138 (1991) 88. 23 D.M. Shub, M.F. Reznik and V.V. Shalaginov, Elektrokhimiya, 21 (1985) 937. 24 S. Ardizzbne, G. Fregonara and S. Trasatti, Electrochim. Acta, 35 (1990) 263. 25 C. Angelinetta, M. Falciola and S. Trasatti, J. Electroanal. Chem., 214 (1986) 535. 26 B.E. Conway, L. Bai and M.A. Sattar, Int. J. Hydrogen Energy, 12 (1987) 607.
.I. Electroanal. Chem., 339 (1992) 255-268
Elsevier Sequoia !&A.,Lausanne
JEC 02074
Hydrogen evolution on oxide electrodes: Co,O, in alkaline solution * E. Veggetti,
I.M. Kodintsev
l
* and S. Trasatti
l
**
Department of Physical Chemistry and Electrochemistry, University of Milan, via Venezian 21, 20133 Milan (Italy)
(Received 11 February 1992)
Abstract
Mechanistic and electrocatalytic aspects of H, evolution from alkaline solutions were investigated using Co,O, cathodes prepared by thermal decomposition of cobalt nitrate. Two different arrangements were used: Ti/Co,O, (A) and Ti/RuO, /C&O, (B). The surface response was investigated by cyclic voltammetry, while the kinetics was studied by determining Tafel lines and reaction orders. The results show that electrodes B are more active than electrodes A with an overpotential decrease of more than 0.2 V. The mechanisms are also different: the reaction order is fractional with electrodes B. The experimental picture for B electrodes is consistent with H, evolution taking place on RuO, at the bottom of pores and/or cracks in the Co,O, layer. Presumably, the latter oxide is cathodically reduced to some form of hydroxide which is not particularly active for H, evolution. In addition, the oxide layer is subjected to mechanical erosion due to the action of the evolving gas.
INTRODUCTION
Oxide electrodes order
to improve
(DSA@) were the performance
introduced into electrochemical technology in electrodes [1,2]. of gas CO,, Cl,) evolution
Ernest Yeager was among the first to investigate their electrocatalytic properties, particularly for 0, evolution and reduction [3]. Recent patents have also claimed good performances for DSA-type electrodes (including RuO,, 110, and spinels) in
Dedicated to Professor Ernest B. Yeager on the occasion of his retirement and in recognition of his contribution to electrochemistry. l * On leave from Mendeleev Institute of Chemical Technology, Moscow, Russian Federation. l ** To whom correspondence should be addressed.
l
0022-0728/92/$05.00
0 1992 - Elsevier Sequoia S.A. All rights reserved
256
the cathodic evolution of hydrogen [4-71. However, there has been little fundamental research in this field. Only a few papers in the open literature deal with the study of hydrogen evolution at oxide electrodes [8-101. The first investigation was carried out in this laboratory long ago [ill before any practical implication was envisaged. Oxides are thermodynamically unstable in the hydrogen adsorption region [12,13] and are thus in principle not expected to be good electrocatalytic cathodes. A prerequisite for understanding the good performances obtained in practice is to elucidate the kinetic mechanism and the surface response. In previous work, we have investigated the electrocatalytic behaviour of IrO, and RuO, cathodes in acid solutions [14-161. In this paper, results of a study of C&O, in alkaline solutions are reported.
EXPERIMENTAL
Electrodes C&O, layers on a Ti support were prepared by thermal decomposition of Co(NO,), .6H,O (Fluka) at temperatures in the range 230-500°C. The precursor was dissolved in isopropanol so as to give a 0.3 mol dmW3 solution and spread onto the support by brushing. After each operation, the solvent was evaporated at 60-70°C and the dried layer was fired for 10 min in a oven preheated to the required temperature. The operation was repeated until the desired oxide loading (as determined by weighing) was reached. Electrodes were finally calcined at the same temperature for 1 h. The typical electrochemical surface properties of Co,O, electrodes prepared in this way have been described elsewhere [17]. The average oxide loading was 3.6 f 0.2 mg Co,O,, corresponding to a nominal thickness of about 3 pm. The mounting of the electrodes and the special Teflon holder have been described elsewhere [18]. Co,O, was prepared at 230, 260, 300, 350, 400, 450 and 500°C. Two samples were prepared at each temperature. Support Ti platelets of dimensions 10 mm X 10 mm X 0.2 mm were used as a support. The metal was first sandblasted and then etched with boiling 10% oxalic acid for 10 min (just prior to depositing the oxide layer) in order to remove any insulating TiO, film. In order to minimize the ohmic drops related to TiO,, which interfere with kinetic determinations, a set of samples fired between 400 and 500°C was prepared by placing an interlayer of RuO, between Ti and Co,O,. A 0.02 mol dmV3 RuCl, * nH,O solution in isopropanol was used to prepare a thin RuO, layer by thermal decomposition at the same temperature.
Solution
Aqueous solutions of 1 mol dme3 NaOH prepared volumetrically were used throughout. They were deaerated before each run and stirred by bubbling nitrogen. The three-compartment cell has been described elsewhere [19]. Experiments were conducted at 25 f O.l”C in a water thermostat. Reference electrode
Potentials were measured and are reported on the saturated calomel electrode (SCE) scale. The reference electrode was connected to the cell through a Luggin capillary to minimize uncompensated ohmic drops. Techniques The surface response was investigated by means of voltammetric curves recorded at 20 mV s- ‘. The kinetics of hydrogen evolution was studied by using quasi-stationary current-potential curves. Currents were read after 3 min at each potential. Potentials were increased in 10 mV steps up to a current of approx. 0.1 A cm-’ and then reversed. Any hysteresis vanished after the second cycle. The order of reaction with respect to OH- was determined by varying the NaOH concentration between 0.1 and 1 mol dme3 while keeping the ionic strength constant with additions of NaClO,. AMEL instrumentation was used throughout. RESULTS
AND DISCUSSION
Voltammetric curves
Voltammetric curves were used to monitor the surface state of the electrodes and to verify the presence of uncompensated ohmic drops related to TiO, films.
Fig. 1. Typical voltammetric curve obtained between -0.1 NaOH for a Ti/Co,O, electrode prepared at 260°C.
and 0.5 V SCE at 20 mV s-’
in 1 mol drne3
258
KXE)
Fig. 2. Voltammetric curve obtained between -0.1 and 0.5 V SCE at 20 mV s-’ in 1 mol dm-3 NaOH for Co30, electrodes prepared at 500°C: (1) Ti/Co,O,; (2) Ti/RuOz /Co,O,.
“Standard” curves were recorded between -0.1 and 0.5 V (just prior to oxygen evolution). Figure 1 shows a typical voltammetric curve for Co,O, in alkali. Details have been given elsewhere [ 171. The curve is characterized by a pair of peaks around 0.4 V related to the Co(III)/cO(IV) redox transition. This surface transition is reversible and therefore the anodic and cathodic peaks should coincide. The separation between the two peaks can be taken as proportional to the uncompensated ohmic drop. Up to 350°C the separation AE between the peaks is small and this indicates the absence of large ohmic effects. At higher temperatures the curves are distorted. However, as Fig. 2 shows, a typical curve is again observed at these temperatures if some RuO, is present in the interlayer. Therefore the distortion is attributed to the insulating properties of TiO,, which is heavily doped by RuO, when an interlayer is present. Voltammetric charge The voltammetric charge q * obtained by integration of the voltammetric curves is determined by the (formal) redox transition [17] CoO(OH)
+ Coo, + H++ e-
(1)
The charge measures the number of surface active sites and can therefore be taken as proportional to the true surface area.
259
8 I
300
I
I
. . 1
400
. I
500
t/C
Fig. 3. Variation of the initial voltammetric charge 4: with the calcination temperature: A Ti/RuO, /Co,O,.
0 Ti/Co304;
Figure 3 shows the variation of q: with the calcination temperature. q* is seen to decrease monotonically as expected for a surface area effect. At T 2 400°C q* drops for the samples without RuO,; however, this is interpreted as an artefact resulting from the distortion of the voltammetric curves. In fact, the qi* values for the RuO,-containing samples at 400°C match the set of values obtained for the other electrodes. Thus a continuous variation of qi* is obtained over the temperature range investigated. Effect of the negative potential limit If the negative limit of the potential range is progressively lowered, the voltam-’ metric curve is dramatically modified. As the potential enters the hydrogen evolution region, a prominent anodic peak develops around - 0.7 V. This is shown in Fig. 4. Comparison with results in the literature obtained with Co(OH), electrodes [20-221 suggests that the onset of atomic hydrogen formation catalyses the reduction of the oxide layer with penetration of hydrogen into the bulk of the oxide. This is an important difference with respect to RuO, and IrO, electrodes which do not exhibit major modifications of the voltammetric curve upon hydrogen evolution 114,161.It is very likely that thermal Co,O, is hydrogenated to hydroxide. The modification is irreversible because the voltammetric charge starts increasing after the imposition of a cathodic load. This aspect will be discussed in detail later on.
260
Fig. 4. Effect of hydrogen evolution on the voltammetric curve at 20 mV s-’ in 1 mol dmm3 NaOH. The electrode was held at - 1.3 V SCE for (1) 10 s, (2) 30 s, (3) 1 min and (4) 5 min.
It is interesting that the electrodes containing RuO, show a much lower tendency towards modification of the voltammetric pattern. It is probable that the presence of RuO, influences also the behaviour of Co,O, as well as that of the support-oxide interface. Current-potential curves Figure 5 shows E-log j curves in the region of hydrogen evolution for two electrodes, one with (electrode B) and the other one without (electrode A) the RuO, interlayer. The difference is dramatic. While hydrogen evolution starts around - 1.05 V SCE for electrodes B, it does not begin until - 1.25 V for electrodes A and is preceded by a slow reaction (which is assumed to be hydrogen penetration into the oxide) with an apparent Tafel slope of about 180 mV. In both cases the main Tafel line is distorted by uncompensated ohmic drop. Analysis of the deviations [23] from the expected straight line reveals that the actual Tafel slope for electrodes without RuO, (A electrodes) is 60 mV if an uncompensated ohmic component of OS-O.6 51 is assumed. This value is typical of ohmic drops in solution at the Luggin capillary. The fact that the same value is obtained at all temperatures suggests that any TiO, layer is presumably reduced on cathodic load and no longer presents a significant ohmic resistance. It is also interesting to note that the calcination temperature does not appear to have any appreciable influence on either the shape or the position of the polarization curve. In other words, surface area effects do not show up at this stage. The Tafel lines of electrodes with RuO, (B electrodes) are less distorted because they start at lower current densities. The Tafel slope at low overpotentials is very close to 40 mV, and is the same for different calcination temperatures. No surface area effect can be detected in this case. This suggests that the active
261
1
I
-4
I
-3
I
I
-1
lo&/A cnY2j2
Fig. 5. Quasi-stationary current-potential curves for H, evolution on Co,O, electrodes: (1) Ti/Co,O,; (2) Ti/RuO* /Co,O,. Tafel lines from ohmic drop corrections are shown. The slopes of the lines are indicated.
surface taking part in the reaction does not depend on the calcination temperature. In the high overpotential region of B electrodes, the analysis of ohmic drop deviations reveals that a second Tafel line with a slope close to 120 mV is present. The ohmic component turns out to be slightly higher with this type of electrode, being about 1.0 fl. Reaction order The current was measured at - 1.24 V for electrodes without RuO, (A) and at - 1.07 V for electrodes with RuO, (B) so as to stay in the linear section of the Tafel line as the pH was changed. Figure 6 clearly shows that the reaction order is definitely different, in the two cases, with a value of - 1 (the reaction rate is depressed by increasing NaOH concentration) for electrodes A and a value of - 0.5 (fractional) for electrodes B. Fractional reaction orders have previously been observed for H, evolution with IrO, and RuO, in acid solutions [14,1.5]. H,-induced modifications H, evolution induces modifications on Co,O, electrodes irrespective of the presence or the absence of .the RuO, interlayer. Immediately after the runs shown in Fig. 4, the “standard” voltammetric curves revealed an increased value of charge. Figure 7 shows that this charge is proportional to the initial charge
262
E >
(4
\ x0
m-
j
\
\
\
\
l’ \
a2t 13
13.4
\
'e
\
\ 13.8
Fig. 6. Dependence of the H, evolution reaction rate at constant potential on the solution pH: (a) Ti/Co,O,, E = - 1.24 V SCE, (b) Ti/RuO, /C&O,, E = - 1.07 V SCE. Straight lines (- - - - - -) of slope (a) 1 and (b) 0.5 are shown.
obtained with fresh electrodes [16]. Thus only subsurface sites are activated by a moderate H, discharge so that a “memory” of the surface area ratio between different samples is still retained. After extensive H, evolution (i.e. after determination of the Tafel line and reaction order) the charge becomes approximately independent of calcination temperature. This is shown in Fig. 8 where the charge after recording the curves in Fig. 4 is plotted against the final charge. Although some scatter is observed, the final value of the charge is about 160 mC crne2. It should be noted that this value is more than five times higher than the highest initial value. Since 1.8 mg cmW2 corresponds to 7.5 pmol cmA2 Co,O,, th e charge associated with (roughly) one electrode exchange per Co atom amounts to about 2 C cmm2. This is approximately 10 times higher than the final charge effectively measured. This may indicate that complete bulk reduction does not take place, and that only the total surface (“outer” + “inner”, i.e. including cracks, pores and grain boundaries), with an approximate ratio of 1: 10 between surface and bulk atoms, is reduced. However, two facts have to be borne in mind.
263
/ 40-
/ l
k
/
/
/
::’
A
20-
/
/
’ k
,$ ,d
i I
10
I
I
20 30 4r/mC cm-*
I
40
Fig. 7. Plot of the voltammetric charge 4: (between -0.1 and 0.5 V SCE) from Fig. 2 for freshly prepared electrodes vs. the voltammetric charge q: after recording the curves in Fig. 4: - - - - - straight line with unit slope; l Ti/Co,O,; A Ti/RuO, /Co30,.
(1) The voltammetric curve is recorded at a finite sweep rate (20 mV s-l). If hydrogen penetration is slow, the diffusion front can penetrate only marginally during a voltammetric cycle [24]. (2) Erosion of the overlayer is visible to the naked eye. The measured final charge could refer to a much smaller thickness while involving reduction of the whole layer. The latter explanation seems the more probable on the basis of the complete experimental picture. If the final charge still involved only surface atoms, it would be difficult to understand how it can be independent of the calcination temperature. Since the particle size differs at different temperatures, a difference in surface area must persist even after extensive surface wetting promoted by H, evolution. This particular aspect deserves further experimental investigation. However, it is also possible that complete “hydrogenation” of the C&O, is a very slow process, and
264
“i .
JpJo- ‘.I
.
z ... .: l
\‘6
. .
.
r.., 50
100 c@nC cn?
150
Fig. 8. Plot of q: (from Fig. 7, same symbols) vs. the voltammetric charge after extensive H, evolution (q%)(at the end of all experiments).
what is seen in this work is only its beginning; nevertheless, it is already sufficient to smooth down surface area effects. It is intriguing that a higher value of the final charge is on average obtained for electrodes with an RuO, interlayer. This is probably due to the fact that the measured charge includes that exchanged by the surface of RuO, itself. Electrocatalytic activity It has already been mentioned that the rate of hydrogen evolution does not appear to depend on the calcination temperature. This is clearly shown in Fig. 9 where the apparent current density is plotted versus the final charge in a log-log
0.5 9 : E > j
O-
-0.5l 1.5,
I
I
2.5 log (s’/mC cm? Fig. 9. Log-log plot of the reaction rate for H, evolution against the final voltammetric charge: - - - - - iine with unit slope. Symbols as in Fig. 7. 2D
265
a4-
e,3-
E
A \
> !? 2
\*
0.2-
\ A \ A
0.1 -
\A
\
\
0
-‘-$--,--g--.
1
0
_a---____ *
I
I
I
300
400
!500
t/‘C Fig. 10. Dependence of the reaction rate for H, evolution normalized to unit surface charge on the calcination temperature: (1) Ti/Co,O.,; (2) Ti/RuO, /C&O,.
diagram. Undoubtedly, there is no correlation between j and qf*; the former remains on average constant as the latter changes. More specifically, Fig. 10 shows a plot of the current density per unit surface charge as a function of the calcination temperature. The electrocatalytic activity per active surface site is seen not to vary in any systematic way with the calcination temperature. Kinetic mechanism The results of the kinetic studies are summarized in Table 1. The reaction mechanism is definitely different for the two sets of electrodes. The 60 mV Tafel slope indicates that the rate-determining step is a chemical reaction following the primary discharge. However, the 40 mV Tafel slope suggests that the second
TABLE 1 Experimental kinetic parameters for H, evolution on Co,O, in alkaline solutions Electrode type
Tafel slope b /mV
Reaction order voH-
Ti/CosO., (A) Ti/RuO, /Co,O., (B)
approx. 60 mV approx. 40 mV
approx. - 1 approx. - 0.5
266
electron transfer is rate determining. In the absence of more detailed molecular information, a common mechanism with two different rate-determining steps can be envisaged:
I
-M-OH
+ H,O + e-
I
-M-OH,*
I
+ e-
I
-M-H+H,O
-M-OH:
+ OH-(aq)
(2a)
I
-
-M-OH,
I
-
-M-OH,
P) I
-
-M-H
I
-
-M-O
PC)
+ OH-(aq)
I
+ H, t
(2d)
The asterisk indicates an unstable surface species which needs to rearrange to a more stable species to become active for H, evolution. Results indicate that H, evolution on Ti/Co,O, electrodes (A) is preceded by an anomalous Tafel region which might describe the slow primary reduction of surface sites (b = 120 mV> with a transfer coefficient lower than 0.5 (based on the assumption of a dual barrier model). H, evolution on these electrodes is probably limited by the slow rearrangement of the reduced surface sites, i.e. step (2b) (b = 60 mV), so that steps (2~) and (2d) only take place at much higher overpotentials. With Ti/RuOJCo,O, electrodes both step (2a) and (2b) are fast so that step (2~) becomes rate determining (b = 40 mV). The proposed mechanism is also consistent with the observed reaction order. Previous results [25] have shown that the surface charge of oxides is pH dependent so that the electric potential at the reaction site varies by 59 mV per pH unit: $J* a - (RT/F)
ln[OH-]
If step (2b) is rate determining,
(3) the kinetic equation is
ja[-OH;]
(4)
Since step (2a) is in quasi-equilibrium, [-OH:]
a [OH-]-’
exp( -FE/Z?T)
(5)
From eqns. (4) and (5) j a [OH-]-’
exp( -FE/RT)
which reproduces
the kinetic parameters
(6) in Table 1 for Ti/Co,O,
electrodes.
267
If step (24 is rate determining,
the kinetic equation is
ja[-OH,]exp[-cy,(E+4*)F/RT]
l
(7)
Since [-OH,]
a [OH-]-’
exp( -FE/RT)
(8)
it follows from eqns. (31, (7) and (8) that j
a
[OH-]-“’
exp[ - (1 + cz,)FE/Z?T]
(9) where ac is the cathodic transfer coefficient. Assuming (Y,= 0.5, eqn. (9) reproduces the kinetic parameters in Table 1 for Ti/RuO,/Co,O, electrodes. The Tafel slope close to 120 mV at higher overpotential for B electrodes is probably related to a shift of the rate-determining step from step (2~) to step (2a) (primary discharge). This is quite often observed in hydrogen electrocatalysis mechanisms [ 10,261. In view of the similarity of the voltammetric curves for electrodes with and without an RuO, interlayer, which indicates that the chemical nature of the surface sites is unchanged, the only possible explanation for the different kinetic mechanisms is that H, evolution takes place on the RuO, sites at the bottom of pores, cracks and crevices. A similar situation has been observed with RuO, layers covered with chemically deposited Ni [lo]. This would account for the additional ohmic loss along the pores which is not observed for electrodes without RuO,. CONCLUSIONS The present results show that Ti/Co,O, electrodes are not very active for the cathodic H, evolution reaction in alkaline solution. Under a cathodic load Co,O, is reduced and H, probably penetrates into the oxide bulk. H, evolution takes place on the reduced oxide with a higher overpotential (> 0.2 V) than expected. The reaction rate is limited by a chemical step following the primary discharge, probably involving some rearrangement of the surface intermediate (a reduced surface species). Ti/RuO,/Co,O, electrodes are more active. H, evolution takes place in the expected potential range so that the overpotential gain is > 0.2 V with respect to pure C&O,. The active surface is not that of Co,O, but that of RuO, at the bottom of pores, cracks, loose grain boundaries etc. The kinetic mechanism is different, as the reaction rate is limited by the transfer of the second electron (further reduction of surface intermediates formed by reduction of water molecules). The observed fractional reaction order indicates that a surface charging effect, related to the acid-base properties of the oxide, is operative.
l In the case of metals, double-layer corrections are written as E - c$* since q~* decreases as the ionic concentration increases. In the case of oxides the correction must be written as E + 4* since the pH governs the surface charge and S#J*increases as the concentration of the potential-determining ion increases.
268
The intense evolution of H, gas provokes erosion of the active layer. For this reason the final charge associated with the redox transition of Co,O, is far less than that calculated for the initial catalyst loading. It thus appears that pure Co,O, does not show particularly interesting electrocatalytic properties for H, evolution in alkaline solution, although it may be useful in specific cases. ACKNOWLEDGEMENT
Financial support of this work by the CNR (National Research Council, Rome) is gratefully acknowledged. REFERENCES 1 S. Trasatti (Ed.), Electrodes of Conductive Metallic Oxides, Parts A and B, Elsevier, Amsterdam, 1980, 1981. 2 S. Trasatti and W.E. O’Grady, Adv. Electrochem. Electrochem. Eng., 13 (1980) 177. 3 W.E. O’Grady,‘C. Iwakura, J. Huang and E. Yeager, in M.W. Breiter (Ed.), Electrocatalysis, Electrochemical Society, Pennington, NJ, 1977, p. 286. 4 E. Nicolas, Eur. Pat. Appl. 23368, 1981; Chem. Abstr., 94 (1981) 199838. 5 J. Clerc-Renaud, F. Leroux and D. Ravier, Eur. Pat. Appl., EP 240413, 1987; Chem. Abstr., 108 (1988) 45864. 6 J.F. Cairns, D.A. Denton and P.A. Izard, Eur. Pat. Appl., EP 129374, 1984; Chem. Abstr., 102 (1985) 102442. 7 A. Nidola, PCT Int. Appl., WO 86 03790 (19861; Chem. Abstr., 105 (1986) 122974. 8 E.R. Kijtz and S. Stucki, J. Appl. Electrochem., 17 (1987) 1190. 9 M. Jaccaud, F. Leroux and J.C. Millet, Mater. Chem. Phys., 22 (1989) 105. 10 A. Anani, Z. Mao, S. Srinivasan and A.J. Appleby, J. Appl. Electrochem., 21 (1991) 683. 11 D. Galizzioli, F. Tantardini and S. Trasatti, J. Appl. Electrochem., 5 (1975) 203. 12 M. Pourbaix, Atlas d’Equilibres Electrochimiques a 25°C Gauthiers-Villars, Paris, 1963. 13 S. Trasatti and G. Lodi in S. Trasatti (Ed.), Electrodes of Conductive Metallic Oxides, Part A, Elsevier, Amsterdam, 1980, p. 301. 14 J.F.C. Boodts and S. Trasatti, J. Appl. Electrochem., 19 (1989) 255. 15 J.F.C. Boodts, G. Fregonara and S. Trasatti in F. Hine, B.V. Tilak, J.M. Fenton and J.D. Lisius (Eds.), Performance of Electrodes for Industrial Electrochemical Processes, Electrochemical Society, Pennington, NJ, 1989, p. 135. 16 S. Ardizzone, G. Fregonara and S. Trasatti, J. Electroanal. Chem., 266 (1989) 191. 17 R. Boggio, A. Carugati and S. Trasatti, J. Appl. Electrochem., 17 (1987) 828. 18 R. Garavaglia, C.M. Mari and S. Trasatti, Surf. Technol., 23 (1984) 41. 19 G. Lodi, E. Sivieri, A. De Battisti and S. Trasatti, J. Appl. Electrochem., 8 (1978) 135. 20 L.D. Burke, M.E. Lyons and O.J. Murphy, J. Electroanal. Chem., 132 (1982) 247. 21 H. Gomes Meier, J.R. Vilche and A.R. Arvia, J. Electroanal. Chem., 138 (1982) 367. 22 L.D. Burke and M.M. Murphy, J. Electrochem. Sot., 138 (1991) 88. 23 D.M. Shub, M.F. Reznik and V.V. Shalaginov, Elektrokhimiya, 21 (1985) 937. 24 S. Ardizzbne, G. Fregonara and S. Trasatti, Electrochim. Acta, 35 (1990) 263. 25 C. Angelinetta, M. Falciola and S. Trasatti, J. Electroanal. Chem., 214 (1986) 535. 26 B.E. Conway, L. Bai and M.A. Sattar, Int. J. Hydrogen Energy, 12 (1987) 607.