Catalytic decomposition of hydrogen peroxide in alkaline solutions

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Thursday Nov 11 01:05 PM StyleTag -- Journal: ELECOM (Electrochemistry Communications) Article: 136

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Electrochemistry Communications 1 (1999) 614–617

Catalytic decomposition of hydrogen peroxide in alkaline solutions

Rajeev Venkatachalapathy a, Guadalupe P. Davila b, Jai Prakash a,*a Center for Electrochemical Science and Engineering, Department of Chemical and Environmental Engineering, Illinois Institute of Technology,

10 W 33rd Street, Chicago, IL 60616, USAb Mechanical Engineering Department, University of Texas-Pan American, 1201 West University Drive, Edinburg, TX 78539, USA

Received 15 September 1999; received in revised form 1 October 1999; accepted 4 October 1999

Abstract

Catalytic activity of carbon, platinum-supported on high-area carbon, platinum, lead ruthenate, and ruthenium oxide towards hydrogenperoxide decomposition in alkaline solution is investigated using the rotating disk electrode technique. The heterogeneous rate constant forperoxide decomposition on these catalysts is determined from the slope of log(iL) versus time, where iL is the diffusion-limiting currentcorresponding to the concentration of peroxide at a given time. The order of catalytic activity is found to be platinum)lead ruthenate)ruthenium oxide)carbon. A general reaction mechanism for the peroxide decomposition on these catalysts is also proposed. q1999Elsevier Science S.A. All rights reserved.

Keywords: Peroxide decomposition; Ruthenium pyrochlore; Lead ruthenate; Heterogeneous rate constant; Rotating disk electrode

1. Introduction

In an effort to achieve an acceptable performance of fuelcells, it has become necessary to resort to expensive precious-metal oxygen electrocatalysts in very high-area forms, whichare difficult to maintain over long periods of time. In addition,due to low solubility in concentrated alkali, the transport ofoxygen by diffusion through the solution phase is very slow.Consequently, it becomes necessary to use porous electrodeswith gas-filled channels, which provide relatively short pathsfor oxygen to reach the reaction site on the electrode surface.These porous electrodes are fabricated with high-area carbonas a conductive support and a catalyst for the oxygen reduc-tion or peroxide decomposition. The carbon surface used inthese electrodes is a catalyst in alkaline solution in its ownright and reduces O2 to HO2

y by a two-electron reductionprocess:

y y yO qH Oq2e ™HO qOH (1)2 2 2

The exchange current densities and the carbon area aresufficiently high that the electrode potential prevailing locallywithin the porous electrode in alkaline electrolytes isexpected to follow the Nernst equation for the two-electronreduction of oxygen to peroxide (Eq. (1)). In order todecrease the cathodic polarization and to enhance the elec-

* Corresponding author. Tel.: q1-312-567-3639; fax: q1-312-567-8874;e-mail: [email protected]

trode life, it is critical to eliminate HO2y formed, according

to Eq. (1). The peroxide activity can be depressed to a lowvalue by the use of an additional catalyst, which can eliminateperoxide effectively as shown below:

y yHO ™OH q1/2 O (2)2 2

A variety of methods such as gasometric [1], open-circuitpotential-decay method [2,3], and steady-state polarizationusing gas-fed electrodes have been used to investigate thisreaction on Ni ferrites [4], perovskite oxides [5], and metalporphyrins [6]. Most of these studies explained the homo-geneous and heterogeneous decomposition of H2O2 on thesecatalysts by the Haber–Weiss mechanism [7], whichinvolves free-radical chain reactions. This reaction, however,does not consider the potential dependence of this reaction.This paper presents the kinetic and mechanistic aspects ofperoxide decomposition on carbon, platinum-supported onhigh-area carbon, ruthenium oxide, and lead ruthenate forreaction 2 using the rotating disk electrode technique.

2. Experimental procedures

Ruthenium dioxide (Aldrich), platinum, platinum sup-ported on carbon XC-72R, 20 wt.% loading (Alfa Aesar),and Carbon XC-72 (Alfa Aesar) were used as received. Thelead ruthenate pyrochlores were synthesized and purified

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Fig. 1. Current vs. potential for a rotating gold disk electrode with 5 mgPb2Ru2O6.5 (3008C heat-treated) dispersed in 70 ml of 1 M KOH. Initialperoxide concentration: 4.29=10y4 M; rotation rate: 60 rps; scan rate: 20mV sy1; disk area: 0.45 cm2; temperature: 258C.

Fig. 2. Logarithm of the limiting current density iL vs. time for differentamounts of lead ruthenate pyrochlore catalyst. Initial hydrogen peroxideconcentration: 4.29=10y4 M; weight of the catalyst is indicated on theplots; rotation rate: 60 rps; temperature: 258C.

according to the method reported by Horowitz et al. [8]. Thisinvolves the reaction of the appropriate metal cations by pre-cipitation and subsequent crystallization of the precipitate ina liquid alkaline medium (4 M KOH) in the presence of O2

at 75–908C. The salts used were Ru(NO)(NO3)3, in aqueoussolution, 1.5% (w/v) (Strem), and Pb(CH3COO)2P3H2O(MCB, reagent grade). The salts were mixed in the appro-priate amounts to achieve the desired stoichiometries andstirred for 20 min. To this solution, was added 4 M KOH ata temperature of 75–808C. Oxygen was bubbled through thesystem continuously. The pH of the reaction mixture wasadjusted between 12.0 and 14.0. The reaction was carried outfor 24–72 h, until the supernatant solution was free of detect-able amounts of the reactant metal ions. The suspension wasallowed to settle overnight and the supernatant solution wasremoved and replaced with distilled, ultrapure water and gla-cial acetic acid in order to leach out any lead oxide formedduring the reaction. The black suspension was then filteredusing a 1 mm Nucleopore membrane filter paper, which wassupported on a Buchner-style funnel. The material was thenheat-treated at high temperatures (300–7508C) in air.

Kinetic measurements of the peroxide decomposition wereperformed by using the rotating disk electrode technique [6].This method is described as follows: a known amount of theoxide powder was added to 70 ml of an Ar-saturated solutionof KOH, and was dispersed by vigorous agitation using amagnetic stirrer. Typically, 1.0 ml of 10y2 M H2O2 (reagentgrade H2O2, FMC) was added to the solution and peroxideconcentration was examined as a function of time. A rotator(PINE Instrument Company), with a control box to adjustthe rotation rate of the electrode, was used to set the rotationrate of the disk electrode to 3600 rpm in order to maintainmass transfer controlled conditions [6]. The decreasing con-centration of the HO2

y was monitored by measuring thediffusion limiting current using a potentiostat (PINE Instru-ment Company). Before every measurement, the Au diskwas polarized at y1.0 V versus HgNHgO for 3 min to removethe oxide layer formed on it. The concentration–time depend-ence was measured at 30 s intervals by scanning the potentialfrom y0.10 to q0.20 V versus HgNHgO at 20 mV sy1. TheKOH solutions were prepared from low carbonate KOH pel-lets (Fluka, puriss., p.a.) dissolved in distilled reverse osmo-sis water.

3. Results and discussion

Fig. 1 shows a typical relationship between the disk cur-rents and potential measured at 60 rps at various time inter-vals, in 70 ml of 1 M KOH containing the catalyst (5 mg ofPb2Ru2O6.5 in the present study). The decrease in the disklimiting currents with time in Fig. 1 is due to the decompo-sition of peroxide. Plots of log iL versus t for differentamountsof Pb2Ru2O6.5 are shown in Fig. 2. The observed linear rela-tionship in this figure indicates the first-order dependence ofthe kinetics of peroxide decomposition on the peroxide con-

centration, which is in good agreement with the literature[4]. The equation governing the first order kinetics is

yd[H O ]2 2 sk[H O ] (3)2 2dt

where k is the apparent first-order rate constant in sy1. Eq.(3) was followed irrespective of the heat-treatment temper-ature of the catalyst or the pH of the solution. The limitingcurrent [9] relates to the changing concentration of peroxideby

U(y1/6) 2/3 1/2i s0.62 nFA n D f C (4)L

where n is the number of electrons, A the electrode area(cm2), and the kinematic viscosity (cm2 sy1) [10],n H O2 2

the diffusion coefficient (cm2 sy1) [10], CU theD H O2 2

concentration of peroxide (M) and f the rotation rate (rps).The apparent rate constants (k) were obtained from theslopesof the plot in Fig. 2. To compare the intrinsic activity of the

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Table 1Rate constant ks for peroxide decomposition on different catalysts using agold RDE

Catalyst BET area/m2 gy1 ks/cm sy1

Platinum 10.0 2.89=10y4

Platinum a 10.0 1.44=10y3

Pb2Ru2O6.5 34.27 1.41=10y4

RuO2 85.87 4.15=10y5

Carbon 250 5.04=10y6

a ks for Pt was calculated based on 20% loading in the Pt/carbon (XC-72R)catalyst and BET surface area of 10 m2 gy1.

Fig. 3. Effect of KOH concentration on hydrogen peroxide decompositionkinetics on lead ruthenate pyrochlore catalyst. Catalyst concentration: 5 mgin 70 ml of 1 M KOH; initial hydrogen peroxide concentration: 4.29=10y4

M; rotation rate: 60 rps; temperature: 258C.

Fig. 4. Dependence of the rate constant k on the activity of the KOH solution.Conditions are the same as in Fig. 3.

different catalysts under various conditions, it is necessary tonormalize the apparent rate constant to the surface area of thecatalyst. Falcon and Carbonio [5] have shown that the het-erogeneous reaction for peroxide decomposition is slow.Under such conditions, the transport to the surface wouldtherefore be diffusion controlled. The heterogeneous rateconstant under these conditions can be expressed as

V kliqy1k (cm s )s (5)s A cat

where Vliq is the volume of solution (in ml), and Acat thesurface area of the catalyst.

The catalytic activity of various catalysts towards peroxidedecomposition is reported in Table 1. The order of the cata-lytic activity is Pt)lead ruthenate)ruthenium dioxide)

carbon. It can also be seen from this table that the catalyticactivity of ruthenium pyrochlore for peroxide decompositionis comparable to that of platinum. This result indicates thatlead ruthenate can be effectively used as a second catalyst onthe cathode side of the alkaline fuel cell to decompose per-oxide, thereby increasing the electrode life and reducing thecathodic polarization.

Fig. 3 shows plots of log(i) versus t for various concen-trations of KOH. The effect of OHy concentration on H2O2

decomposition on Pb2Ru2O6.5 is shown in Fig. 4. In this figurethe activities rather than the concentration of KOH are plottedagainst the corresponding rate constants. The linear relationin Fig. 4 is consistent with the equation:

yln ksym ln[OH ]qA (6)

where m is the proportionality factor and A the intercept.The peroxide decomposition on these catalysts can be

explained by the following reaction scheme involving twoindividual electrochemical steps, which are coupled witheachother:

y y yO qH Oq2e ™OH qHO (7)2 2 2

y y yHO qH Oq2e ™3OH (8)2 2

Reaction 7 is the overall irreversible reaction and it maygo through a sequence of intermediate steps. This mechanismis consistent with the one proposed by Yeager and co-workers[11] to explain peroxide decomposition on a mixed oxide ofmanganese involving two individual electrochemical steps.The current density for H2O2 decomposition can be shown tofollow Eq. (9) at higher negative potential:

yln isA9yk0 ln[OH ] (9)

where

yA9sln i qln[O ]yaF/RT[E (O /HO )o 2 o 2 2

y yyE (HO /OH )] (10)o 2

Since Eq. (9) is in good agreement with Eq. (6), it can beproposed that the H2O2 decomposition on Pb2Ru2O6.5 pro-ceeds by the electrochemical mechanism.

4. Conclusions

The ks values give a clear indication of the order of thecatalytic activity of the different catalysts for peroxidedecomposition. The order of the catalytic activity is Pt)leadruthenate)ruthenium dioxide. The catalytic activity of leadruthenate is comparable to that of platinum. This indicatesthat lead ruthenate can be successfully used as a second cat-alyst on the cathode side of fuel cells to decompose peroxide,

R. Venkatachalapathy et al. / Electrochemistry Communications 1 (1999) 614–617 617


thereby increasing the electrode life and reducing thecathodicpolarization.

Acknowledgements

The authors would like to thank Dr S. Donepudi for thetechnical discussions.

References

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[9] A.J. Bard, L.R Faulkner, Electrochemical Methods, Fundamentals andApplications, Wiley, New York, 1980, p. 288.

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Catalytic decomposition of hydrogen peroxide in alkaline solutions