Fractional reaction orders in oxygen evolution from acidic solutions at ruthenium oxide anodes

Materials Chemistry 6 (1981) 255- 266

FRACTIONAL REACTION ORDERS IN OXYGEN EVOLUTION FROM ACIDIC SOLUTIONS AT RUTHENIUM OXIDE ANODES*

A. CARUGATI**, G. LODI***, S. TRASATTI Laboratory o f Electrochemistry, the University, Via Venezian 21, 20133 M I L A N Italy. Received 22 April 1981; accepted 15 May 1981 Oxygen evolution at RuO2 anodes from acidic solutions exhibits a + fractional reaction order with respect to H . This is close to - 1 for a set of "compact" film electrodes, and close to - 2 for a set of "cracked" film electrodes. The difference in mechanism is related to the different defect structure of the surface. Possible reasons for the observed fractional reaction orders are discussed. It is suggested that the experimental observation may be explained in terms of the reaction proceeding with different mechanisms in parallel on different patches o f the surface whose structure is governed by the acid-base properties of the oxide. Abstract

INTRODUCTION In a previous work I we showed that oxygen evolution takes place on RuO2 anodes with different mechanisms depending on the detailed surface structure and morphology of the active layer. In particular, so-called "compact" (low-defect) electrodes exhibited an average Tafel slope (b) close to 0.04 V increasing with in-. * Paper presented at the 32 ° ISE Meeting, 14-18 September 1981, Dubrovnik, Yugoslavia. Title and abstract submitted in January 1 9 8 1 . ** Presently, holder o f a De Nora fellowship in this laboratory. *** Chemical Institute, the University, Ferrara, Italy. 0390-6035/81/050255 - 1252.00/0 Copyright © 1981 by CENFOR S.R.L. All fights of reproduction in any form reserved

256

creasing compactness, whereas in contrast a b value of about 0.03 V was observed with so-called "cracked" (highly defective) oxides. The proposed mechanism was that known 2 as "electrochemical oxide path" for "compact" f'dms, and that known as "oxide path" for "cracked" films. The main criterion to distinguish between the various possible paths was essentially based on the value of the Tafel slope. Since the two proposed paths differ also in the pH dependence of overpotential, the determination of the reaction order with respect to H + (in acidic solutions) and OH" (in alkaline solutions) comes out to be a powerful tool to discriminate between the above mechanisms. Exhaustive results are now available in this direction and they are shown in this paper.

EXPERIMENTAL "Compact" f'dms were prepared after Lodi et al. z. The temperature of preparation was varied between 300 and 600°C. The films were characterized 4 by measuring the voltammetric charge q* spent to change the potential (at 50 mV s"l) between 0.4 and 1.2 V (the). The charge was measured before and after the kinetic experiments. Oxygen evolution was studied at 25°C by the potentiostatic technique. The base electrolyte was HC104 whose concentration was varied between 0.06 and 1 mol dm "s . Triple distilled water, the second disti[lation being made from alkaline permanganate, was used to prepare solutions. These were stirred in the cell by a stream of oxygen and by a magnetic stirrer. Methods, assemblage and procedure have been described elsewhere l, s, 6. Although the experimental results with cracked electrodes were obtained years before those with compact electrodes, the design of the cell was the same. The Luggin capillary was placed as close as possible to the electrode surface in a side position with respect to the ~knes of current so as to minimize possible ohmic drops. The possible presence of significant IR drop was checked by means of current pulses. It was found to be quite negl~ible over the 2-3 decades of current over which the Tafel slope was determined (maximum value 1 mA cm'2). The constancy in Tafel slope over the concentration range explored (see later, Table 1) provides additional (indirect) evidence for the irrelevance of the ohmic drop correction under the present conditions. The preparation of "cracked" electrodes has been described in previous'papers s, 6. The results reported here for these electrodes in solution of pH =0.5 belong to the same set of data as those in the previous work 1. All other results have

257 not been shown before. Films on quartz were also prepared to investigate the possible influence of the nature of the support. It has been shown previously 1 that the ohmic drop along the f'drn is negligible under the experimental conditions chosen.

RESULTS AND DISCUSSION Table 1 summarizes Tafel slopes, b, and charge, q*, for the various electrodes. In the set of "compact" electrodes, the layer prepared at 300"C on Ti appears to exhibit definitely a lower Tafel slope (ca. 35 mV) than the rest of the group (ca. 40-47 mV). There is a distinct trend of b to increase as q* decreases. This confirms previous observations 1. On the basis of the Tafel slope value, the "electrochemical oxide path" with the second electron transfer being the slow step, is suggested to be the mechanism for oxygen evolution at low-q* electrodes, whereas the "oxide path" 2 with the chemical step being slow, is very likely to be followed at the high-q* electrode. The values of the Tafel slopes do not coincide exactly with those theoretically expected (0.03 and 0.04 V) so that an additional parameter is required to confirm the above suggestion. In the group of "cracked" electrodes, the Tafel slope at pH = 0.5 was found to be equal to 0.03 V for the layers prepared at 300 and 400"C, but an increase to 0.033 V was observed for the f'flrn at 500"C. On the basis of the Tafel slope alone, the most probable mechanism is the "oxide path" with the OH-OH recombination reaction being rate determining. Fig. 1 shows a plot of log j as a function of log CH+. In the concentration range explored, alog CH+/~pH does not differ detectably from 1 7. The solid lines in the main diagram mark the slopes equal to 1 and 2, respectively. The points for all electrodes except electrode No. 1 have been normalized by superposing the various plots. No evidence has been found for a systematic change in slope with the temperature of preparation. The experimental reaction order is close to - 2 for electrode No. 1, and about -1.25 for all other electrodes. Apparently smaller deviations from the integral reaction order could probably be observed to a first approximation if the solid lines were directly superposed to the experimental points. However, in Fig. 1 the straight lines have been drawn starting from the point for the highest H ÷ concentration. In the inset, it can be seen that the reaction order of ca. - 2 is also exhibited

258

!

!

i

a

0 1 '1 2

e

-1.5

/

-1

I

1

~

I

0

!

-I;1'5

0

log [c/tool dm-3] Fig. 1 - Log-log plot for current density for oxygen evolution on R u 0 2 anoae., at E = 1.13 V (sce J against acid concentration in solution. Main diagram: "compact" flow-defect) layers prepared at (A) 300°C on Ti, ( e ) 300°C on quartz, (&) 400°C on 2~, ( n ) 500°C on 2~, ( v ) 500°C on quartz, and ( 0 ) 600°C on quartz. Inset: "'cracked" [highly defective) layers prepared at (A) 400°C, 1/2 t~m, ($) 400°C, 2 I~m, and ( v ) 500°C on 23" C ) Straight lines with 1 and 2 slope. C~rrent on a conventional scale because the different plots [solid symbols) have been superposed by sliding parallel to the ordinate axis.

259 by "cracked" electrodes for which however only two experimental points (pH = 0.5 and 1.9) are available. The limited data for "cracked" electrodes are due to the fact that the behaviour of this group of electrodes has now been revisited in the light of new develop. ments. The earlier experiments were not programmed precisely for the purpose with which they are presented here. However, it should be stressed that the inset is not in fact a two.point plot but rather a six-point plot since the same behaviour was observed with three different "cracked" electrodes. Therefore, despite the paucity of the experimental data, the significance of the plot in the inset should not be minimized. There is enough evidence to rule out that the reaction order is equal to that of "compact" electrodes and the measured value compares well with that expected for a mechanism predicting 0.03 V for the Tafel line. Thus, the previous suggestion i that high values of q* are indicative of high surface concentration of active sites is corroborated and the relative mechanism for oxygen evolution based on a retarded chemical recombination of adsorbed OH groups is confirmed. The results with low-q* electrodes are more complex. The reaction order is definitely higher than unity and the linearity of the plot is not unambiguously ascertainable. It might also be that a curve rather than a straight line fits better to the points. If this is the case, the reaction order ensues to be pH dependent. Despite the fractional value of the reaction order, these results confirm that the mechanism on low-q* electrodes differs from that for high.q* electrodes. The "electrochemical oxide path" with the second electron transfer being the r.d.s., as proposed previously 1, is the more probable mechanisms. Next possibility would be a slow chemical step OHad ~ Oad + H ÷ following the primary discharge of water molecules, but the expected Tafel slope s (0.06 V) is def'mitely higher than the observed one. Fractional reaction orders have been observed also by other authors in acidic solutions. Kokoulina et al.9 have reported that oxygen evolution on TiO2 + RuO2 (30 tool ~) takes place with an almost pH independent overpotential in the low current density region. Actually, some change in overpotential can be seen especially for neutral solutions, but at very high and very low pH's overpotential appears to be effectively pH-invariant. This, with the reported average Tafel slope of 40 mV, gives ca. -1.5 for the reaction order with respect to H ÷ (or 1.5 with respect to OH'). The experimental points are however too much spaced over the explored pH range so that no definite quantitative conclusions can be drawn. Brun~ et al.l o have studied the oxygen evolution at TiO2 + RuO2 (30 tool ~) as a side reaction during chlorine liberation from concentrated brine. They

260 varied the pH over the range 0.5 to 3.5 and their experimental points are so little spaced as to leave no doubt that the reaction order with respect to H ÷ in the pH range 0.5 to 2 is fractional and approximately equal to -1.4. Although the fractional reaction order might be also related in this work to the presence of specifically adsorbed CI" ions, the above results and the results of our work altogether point to the concrete possibility of a fractional reaction order in the evolution of oxygen from acidic solutions. A reaction order of -1.8 with respect to H ÷ has been reported for oxygen evolution on Ru "metal" from acidic solutions 11. The reaction is thought to occur on a hydrated layer of anodically grown RuO 2 12 Fractional reaction orders have apparently not been observed in the case of alkaline solutions. Kokoulina et al.9 have actually suggested that the fractional reaction order extends all over the pH range, but this cannot be taken too seri. ously in view of the low selectivity of their experimental points. In contrast, resuits by O'Grady et al.13 and Wolf et al.14 indicate that the reaction order with respect to OH" is 1 in the pH range 11-14, although the two sets o f experiments differ in the value of the Tafel slope, which is reported to be 0.04 V by O'Grady et al.ll and 0.06 V by Wolf et al.14. The decrease o r b from 0.06 to 0.04 V as the support (Ti) is pre.oxidized as reported by Wolf et al.l ,t is puzzling and in contrast with expectations I s. Possible reasons for the fractional reaction order have never been discussed. Bun6 et al.l o have proposed a mechanism predicting unit reaction order thus neglecting this aspect at all. Kokoulina et al.9 have stated that one cannot make a choice between the various mechanisms satisfying the observed kinetic parameters. They have not discussed the meaning of the fractional reaction order, but have pointed out that the integral reaction order as observed by O'Grady et al.l 3 may be due to the use of poorly buffered solutions. In what follows a number of possible reasons for the observed fractional reaction order is examined. (i) The true surface area changes with the use. Table 1 shows that the value of q*, which can be taken I to be a measure of the surface concentration of active sites is seen to become consistently higher after the evolution of oxygen, with only one exception. This may be an indication o f some sloughing In the surface of electrode No. 1, and of some change in stoichiometry for incorporation of oxygenated species 16 in the other cases. Thus, the fractional reaction order could be an artifact related to changes in the surface concentration of active sites. However, the experimental outcome should in this case depend closely on the procedure of performing experiments. These were carried out in the sequential order 1, 0.25, 0.1, 0.15, 0.06,

261 Table 1 - Parameters for the oxygen evolution reaction on RuO2 electrodes. Electrode a

Support

Tafel slope/V b

" C o m p a c t " layers

Charge q*/mC cm "2 c Initial

Final

1 (300)

Ti

0,035

12.8

8.3

2 (300)

quartz

0.040

4.9

5.4

3 (400)

Ti

0.044

3.5

3.6

4 (500)

Ti

0.046

1.4

1.9

5 (500)

quartz

0.045

2.5

2.9

6 (600)

quartz

0.048

1.8

1.9

7 (400)

Ti

0.030

167

178

8 (400)

Ti

0.030

29

28

9 (500)

Ti

0.033

45

46

" C r a c k e d " layers

a Firing temperature in brackets. The nominal thickness is 2/~m except 0.5 ~m for electrode No. 8. b Mean vah~e to + 0.001 V over six concentrations between 0.06 and 1 tool drn"3 HCIO4 for "compact" electrodes. The values for "cracked" electrodes refer to the solution at pH = 0.5. At pH = 1.9 the Tafel slope increases a little so that the average value over the explored concentration range (at constant ionic strength) is 0.034, 0.034 and 0.037, respectively. c Voltammetric charge between 0.4 and 1.2 V (the) before and after the kinetic experiments. 0.4 mol dm "3 HC104. The increase in q*, in the case o f " c o m p a c t " electrodes,

might partly account for only a small increase in reaction order. Moreover the scat. ter o f points does not follow systematically the sequence o f the experiments which indicates that there is no linear increase in q* with time. The same is also the case with " c r a c k e d " electrodes. This means that the change in q* probably occurs during the first stages o f oxygen evolution experiments. ( i i ) The interfacial pH differs from the solution bulk pH. Since 0 2 evolution is accompanied by liberation o f H ÷, the local concentration o f protons may rise over that in the bulk o f the solution as reaction rate is much higher than diffusion

262

into the liquid phase. This problem has been discussed by Krishtalik 17 and Barral et al. 1s. However, calculations show that in the pH range 0 to 2 as used in this work, the interracial pH is expected not to differ appreciably from the bulk solution pH. On the other hand, the decrease in interfacial pH with respect to the bulk value should diminish rather than raise the apparent reaction order with respect to

H +" (iii) Effect of the diffuse layer potential 19. Experiments have not been carried out at constant ionic strength. Thus, in principle, the observed reaction order might not correspond to constant (E - 02 ), where 02 is the diffuse layer potential. However, also in this case, the reaction order is expected to decrease, as discussed by Conway and Salomon 2°. (iv) Change in reaction mechanism. Transition from a mechanism with one def'mite rate determining step to a mechanism with two parallel reaction paths as the pH is increased may produce an apparent change in.the observed reaction order. Table 1 shows that the Tafel slope does not change with pH for the group of "compact" electrodes, whereas a small increase in Tafel slope is actually observed with "cracked" electrodes. The increase in reaction order with increasing pH may imply a transition from a path whose reaction order is 1 to a path whose reaction order is 2, both exhibiting the same Tafel slope. (v) Acid-base properties of the surface of oxides. Unlike the case of bare metal surfaces, oxides are known to interact strongly with the solution through the surface hydroxyl groups. The state of charge of the oxide surface is determined by its acid-base properties 21 . Thus, even at constant ionic strength, the state of charge may change as the pH is changed. In principle, this corresponds to the occurrence of specific adsorption. If partial dissociation of the OH surface groups occurs, the surface of an oxide may be represented as an ensemble of OH and O" groups, the latter coming from the dissociation of OH groups OH -* O" + H ÷

(1)

Equation (1) applies to only a fraction 5 of the surface. Thus, different reaction mechanisms may take place on the fractions 5 and (1 - 8 ) of the surface. Of the five points discussed above, point (v) in connection with point (iv) seems to be the more reasonable to explain the observed fractional reaction orders. For the electrodes exhibiting the reaction order o f - 1 . 2 5 , the mechanism of oxygen evolution is suggested to start with the fast primary discharge of water molect~les

263 H20 ~- OH + H ÷ + e

(2)

and the rapid equilibration of the formed OH groups with the solution 8OH ~-50" + 8H + .

(3)

Thus, the complete reaction (2) at equilibrium may be written as H20 ~ ( 1 - 8 ) OH + 8 0 " + (1 + 8 ) H ÷ + e

(4)

is expected to be a function of the solution pH. It is then suggested that the further oxidation of OH and O" groups proceeds at comp~able rates but through different pathways. Thus, the fraction (1 - ~ ) of surface OH groups may be oxidized following the "electrochemical oxide" path 2

(! - 8 ) O H - > ( 1 - 8 ) O + ( 1 - 5 ) H + + ( 1 - 5 ) e

(s)

while the remaining fraction'8 of surface O" groups may be oxidized according to Krasil'shchikov path s 80" --, 8 0 + 8 e .

(6)

The detail of the further steps to give 02 as the final product is not important here if reactions (4) and (5) are regarded as controlling the rate in the respective paths. The reaction order with respect to H + is predicted to be -1 on the fraction (1 - 8 ) of surface covered with OH groups and - 2 on the fraction 8 of surface covered with O'. The observed reaction order is thus - (1 - 8) - 28 = - (1 + 8). Since 8 may reasonably be a function of pH, the reaction order may vary with pH. At low pH values 8 should be close to zero because surface dissociation is suppressed by the acidity in solution. Then the limit of the reaction order is possibly - 1 . Fig. 1 shows that a certain curvature might in fact be recognized in the plot and the slope of - 1 is probably more approached at low than at high pH's. For the above reaction mechanism, the observed Tafel slope is given by b = =(1 - 8 ) b s + 8 b 6 , where b s and b 6 are the Tafel slopes when steps (5) and (6), respectively, are rate determining. Since b s = b 6 = 0.04 V, then the observed Tafel slope should be exactly 0.04 V and pH independent. Table 1 shows that this is true only for the electrode prepared at 300°C on quartz, whereas adefinite increase in b

264 is observed as the temperature of preparation of the active layer is increased. Since the reaction order appears not to change with the f'mng temperature, it is possible that this situation is not determined by the value of 8 but rather by a transition from a slow second electron transfer to a retarded primary discharge - reaction (2) - on the fraction (1 - 8 ) of surface covered with OH. In this case b becomes a function of 5 since 2 b 2 = 0.12 V. These results, apart the details of the mechanism that has always relative and not absolute validity, suggest that oxygen evolution on ox. ides is probably a complex reaction as a consequence of the surface being energetically and structurally inhomogeneous. The reaction order close to - 2 with a Tafel slope lower than 0.04 V as observed on "cracked" electrodes and on sample No. 1 of the "compact" electrodes may be explained in terms of the "oxide" path 2 where the reaction 2OH -~ O + H20

(7)

follows reaction (2) and is rate controlling. For this mechanism the predicted Tafel slope is b 7 = 0 . 0 3 V and the reaction order - 2 with respect to H ÷. If it is assumed that equilibrium (3) is also present on these surfaces, then the Tafel slope is expected to be higher than 0.03 V because no other simple reaction step is known 22 with an equal or lower Tafel slope than reaction (7), except O atom recombination which is however hard to envisage as the rate determining step 22. Fig. 1 shows that the reaction order is always less than the theoretically expected - 2 . Thus, a mechanism with two parallel steps may be envisaged, one of which is constituted by reaction (7) and the other may be Krasil'shchikov's path 8 buth with reaction (7) as the rate determining step. In this case the reaction order is predicted to be - 2 (1 - 5 ) - 8 = - ( 2 - 8 ) and the Tafel slope is given by b = = (1 - 5) b 7 + 5b 3 = (1 - 8) 0.03 + 50.06 = 0.03 + 0.035. For "cracked" electrodes the experimental reaction order is ca. - ( 1 . 8 - 1.9); thus, 8 = 0.1 - 0.2 and b = 0.033 - 0.036. In fact, the Tafel slope is higher than 0.03 at pH = 0.5 only for the electrode prepared at 500°C, but it is observed to increase slightly as the pH is increased also with the other electrodes. Thus, the portion of surface involved in reaction (3) probably vanishes as the pH is decreased, and this may be reasonable. A reaction mechanism with steps (6) and (7) proceeding in parallel is not probable since it would predict a reaction order equal to exactly - 2 . A double layer effects as discussed above in (iii) might account for the decrease in reaction order which however should now be z° -1.5. Another possibility is that reaction (5) also proceeds in parallel on the portion (1 - 5 ) of the sur-

265

face. However, the apparently fractional reaction order may also be a consequence of a decrease in the active surface with increasing pH if only the fraction (1 - 5 ) is available for reaction (5) while the portion 5 is inactive (recombination between two charged species of the same sign is highly unlikely). The observed reaction order with sample No. 1 of the "compact" electrodes is about - ( 1 . 7 - 1.8). Thus, 8 = 0.2 - 0 . 3 and b = 0 . 0 3 6 - 0.039 which is not far from the experimental observation. The higher Tafel slope for the "compact" electrode No. 1, as compared to "cracked" electrodes at low pH may reflect a different proportion of the two parallel reactions depending on different features of the surfaces. It is stressed again that the Tafel slope is an average value for "cracked" electrodes because it is pH dependent whereas this is not the case with sample No. 1 of "compact" electrodes. Also in this case, as above, the combination of different steps might satisfy the observed kinetic parameters, but there is no concrete possibility to discriminate in detail between the various possibilities. In conclusion, the above data suggest that the observed fractional reaction orders and the measured Tafel slopes deviating from theoretical values may be understood in terms of parallel kinetic pathways associated with the acid-base properties of the oxide surface. The role of the hydroxylation of the surface in kinetics, never put forward for RuO 2 electrodes before, has been mentioned in a somewhat different way for Co aO 4 electrodes in alkaline solutions by Tarasevich and Efremeov23, who also observed fractional reaction orders in oxygen evolution. Any further quantitative discussion calls for the knowledge of the surface chemistry of RuO2 together with its point of zero charge. No data are presently available in this respect. Work in this direction is in progress in this laboratory:

Acknowledgements The authors are indebted to Dr. G. Buzzanca, now in the laboratories o f C1SE, Segrate (Milan), for the data in the inset in Fig. 1. Financial support from the ATational Research Council ( CNR } is gratefully acknowledged.

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