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Journal of Electroanalytical Chemistry 482 (2000) 40–47

Mechanism of hydrogen peroxide oxidation reaction at a glassycarbon electrode in alkaline solution

P. Westbroek *, E. TemmermanDepartment of Analytical Chemistry, Uni6ersity of Gent, Krijgslaan 281 S12, B-9000 Gent, Belgium

Received 6 October 1999; received in revised form 13 December 1999; accepted 20 December 1999

Abstract

The voltammetric behavior of hydrogen peroxide at a glassy carbon electrode in alkaline solution was studied. An anodicprewave, preceding the well-known transport controlled oxidation reaction of hydrogen peroxide, was observed. The voltammetriccurrent in this wave is independent of the rotation speed of the glassy carbon disc electrode, even in what looks like a limitingcurrent plateau, denoted by a pseudo limiting current. This current is dependent strongly on pH and temperature, does not varylinearly with hydrogen peroxide concentration and the reaction orders, with respect to hydrogen peroxide and hydroxyl ions (pH),are dependent on the hydrogen peroxide concentration. A mechanism that can explain these experimental results has beenproposed. It consists of two submechanisms, occurring simultaneously at the electrode surface, which are in competition with eachother. Their relative predominance is controlled by the hydrogen peroxide concentration, which can explain the varying reactionorders with respect to hydroxyl ions and hydrogen peroxide. The current equation resulting from the proposed mechanism showsthe possibility of using the studied electrode reaction for the amperometric determination of hydrogen peroxide up to highconcentrations. © 2000 Elsevier Science S.A. All rights reserved.

Keywords: Hydrogen peroxide; Glassy carbon; Kinetics; Sensor development

1. Introduction

Sensors for on line measurement of hydrogen perox-ide are of considerable interest for controlling the con-centration of the reagent in various industrialapplications [1–7]. Some of these need a system that isable to measure hydrogen peroxide concentrations upto 2 mol dm−3. Except for a conductometric method[8], there are no electrochemical sensors that can mea-sure such high hydrogen peroxide concentrations. In thecase of voltammetric or amperometric techniques, thisis due to ohmic drop effects. Trials with ultra microelectrodes and membrane covered electrodes to mini-mize ohmic drop did not show promising results be-cause of the limited lifetime and stability of the sensorsignal [9].

In earlier publications the electrochemical behavior ofhydrogen peroxide, reaction mechanisms and analytical

possibilities at several electrode materials were discussed(platinum [10–15], palladium [16], gold [17–19], silver[20–25], carbon [16,21,26] and iron [27–29]). None ofthese papers mentions the occurrence of more than onevoltammetric oxidation wave of hydrogen peroxide.While studying glassy carbon as a possible material foran amperometric sensor electrode for monitoring andcontrolling hydrogen peroxide concentrations, twovoltammetric oxidation waves were observed.

Glassy carbon is a polymer [30] first synthesized byYamada and Sato [31]. A detailed analysis and descrip-tion of the surface was given by Jenkins and Kawamura[32]. Other authors [33,34] reported on the conditionand stability of the glassy carbon surface when it isexposed to air and aqueous solutions for a long periodof time. It is important to use high density glassycarbon to avoid the risk of absorption of oxygen, thereaction product of the oxidation of hydrogen peroxidestudied in this work. High density glassy carbon is amaterial substrate for the determination of the oxygenabsorption characteristics of deposits at other materials[35].

* Corresponding author. Tel.: +32-9-2644820; fax: +32-9-2644992.

E-mail address: philippe.westbroek@rug.ac.be (P. Westbroek)

0022-0728/00/$ - see front matter © 2000 Elsevier Science S.A. All rights reserved.PII: S 0 0 2 2 -0728 (00 )00009 -7

P. Westbroek, E. Temmerman / Journal of Electroanalytical Chemistry 482 (2000) 40–47 41

This paper reports on the first wave, which showspromising characteristics for analytical purposes inhighly concentrated hydrogen peroxide solutions. Thisincludes relatively high oxygen concentrations presentat the solution � electrode surface interface. Thereforethe formation of oxygen gas must be taken into ac-count as a possible interfering parameter in the kinet-ics of the hydrogen peroxide oxidation reaction. Asglassy carbon itself does not absorb oxygen, the com-plexity of the oxygen three phase system (absorbed,dissolved, gas) can be reduced to a two phase system.

2. Experimental

The working electrodes were prepared from glassycarbon rods obtained from Le Carbone Lorraine,type V10. A rod of 3.0 mm diameter and approxi-mately 20 mm length was positioned in a copper baseand embedded in epoxy resin (Epofix Kit fromStruers) [36]. Platinum ring-glassy carbon disc elec-trodes were obtained from Pine Instruments. The sur-face of the electrodes was pretreated in the followingway: to obtain a fresh electrode surface it was flat-tened on SiC paper, type 1200 grit (Buehler). Tosmoothe the relatively rough surface, it was polishedon a polishing cloth (Buehler) with Al2O3 polishingpowder of 1, 0.3 and 0.05 mg granular diameter suc-cessively for 5, 10 and 20 min on a Planopol-2 pol-ishing machine from Struers. To clean the surface ofthe electrode, it was subjected to ultrasonic vibration(Benson, type 3210) for 2 min in a 0.4 mol dm−3

NaClO4-solution of pH 12.0.A carbon rod (Le Carbone Lorraine) and a satu-

rated calomel electrode (Radiometer) served as thecounter and the reference electrode, respectively.

The potentiostat used for the linear sweep andchronoamperometric experiments was typePGSTAT10 from Eco Chemie and was controlled bythe GPES4.4 software package (on an Eknadata486DX2 computer) of the same company. For thering disc experiments, use was made of a PSTAT10in combination with the bipot module, both from EcoChemie.

The pH and temperature measurements were made,respectively, with a combined glass electrode fromSchott and a PT100 temperature sensor from Jumousing a 420A pH-meter from Orion.

All chemicals (hydrogen peroxide, sodium hydrox-ide, sodium perchlorate, sulfuric acid, potassium per-manganate) used in this investigation were ofanalytical grade and obtained from Fluka. Doublydeionized water was obtained from a Millipore-Q-ionexchanger system.

3. Results

Linear sweep voltammetric experiments at a rotat-ing glassy carbon disc electrode in alkaline solutioncontaining 0.4 mol dm−3 NaClO4 were carried out atdifferent hydrogen peroxide concentrations (0.005,0.02, 0.05, 0.1, 0.2, 0.3, 0.4 and 0.5 mol dm−3), pHvalues (10.0, 10.5, 11.0, 11.5, 12.0, 12.5, 13.0, 13.5and 14) and rotation speeds (100, 400, 900, 1600,2500, 3600 and 4900 rpm) of the electrode. An exam-ple of typical current–potential curves is shown inFig. 1. They were recorded at a hydrogen peroxideconcentration of 0.25 mol dm−3, at pH 11.72 and atemperature of 298.0 K, while the rotation rate of theglassy carbon disc electrode was varied, and revealthe presence of two anodic waves. Current–potentialcurves obtained by the variation of pH and hydrogenperoxide concentration are not shown because oftheir similarity with the curves given in Fig. 1. Thecurrent of the first wave, with a half wave potentialof 0.13 V versus SCE, is slightly dependent on rota-tion speed, even in what looks like a limiting currentplateau, further denoted by the pseudo limiting cur-rent. Because of the presence of a large excess ofsupporting electrolyte, migration effects are assumedto be absent. The foregoing indicates that the rate ofthe observed oxidation reaction is not transport con-trolled. Furthermore, the current density of the firstoxidation reaction of hydrogen peroxide is smallerthan the density of the second, transport controlled,oxidation reaction. The first wave is not present whenpH is lower than about ten. In the cathodic part ofthe curve, a reduction wave can be observed corre-sponding to the reduction of hydrogen peroxide andoxygen, with hydroxide ions as the reaction product.The cathodic and the second anodic reaction are notconsidered further in this paper.

In the onset of the first wave, up to a potential ofca. 0.2 V versus SCE, log I varies linearly with poten-tial. Assuming Butler–Volmer behavior, this allowsthe determination of a transfer coefficient if n, thetotal number of electrons transferred, is known. With

Fig. 1. Current potential curves of 0.25 mol dm−3 hydrogen peroxideat a rotating glassy carbon disc electrode (pH 11.72; T=298.0 K) atrotation rates of (1) 100; (2) 400; (3) 900; (4) 1600; (5) 2500; (6) 3600and (7) 4900 rpm.

P. Westbroek, E. Temmerman / Journal of Electroanalytical Chemistry 482 (2000) 40–4742

Table 1Values obtained for the transfer coefficient of the first oxidationreaction of hydrogen peroxide at a glassy carbon electrode in alkalinesolution for several pH values and hydrogen peroxide concentra-tions a

pH c/mol dm−3

0.005 0.02 0.05 0.1 0.2 0.3 0.4 0.5

0.48 0.51 0.4910.0 0.490.47 0.49 0.48 0.510.48 0.47 0.4810.5 0.480.45 0.50 0.52 0.480.48 0.47 0.48 0.500.45 0.4911.0 0.52 0.50

0.4811.5 0.47 0.49 0.50 0.51 0.51 0.51 0.500.5212.0 0.48 0.48 0.48 0.53 0.49 0.52 0.49

0.52 0.46 0.52 0.510.48 0.5012.5 0.48 0.490.47 0.49 0.51 0.48 0.50 0.48 0.4713.0 0.500.49 0.50 0.51 0.490.51 0.4913.5 0.50 0.500.4914.0 0.500.50 0.51 0.50 0.49 0.51 0.52

a T, 298.0 K; N, 900 rpm.

Fig. 2. Logarithmic plot of the relationship between current signal atE=0.45 V versus SCE and hydroxide concentration for the firstoxidation wave of hydrogen peroxide at a glassy carbon electrode;T=298 K. Hydrogen peroxide concentrations are (x) 0.02; (�) 0.05;() 0.10; ( ) 0.20; () 0.30; (�) 0.40 and (�) 0.50 mol dm−3.

The relationship between pseudo limiting current andhydrogen peroxide concentration can also be obtainedfrom the experimental data obtained in the secondseries of experiments. This is presented in Fig. 3 forsome of the pH values investigated (10.0, 10.5, 11.0,11.5, 12.0, 12.5, 13.0, 13.5 and 14). The relation be-tween hydrogen peroxide concentration and pseudolimiting current is nonlinear and deviates from linearitywith increasing pH. The current depends strongly onthe alkalinity of the solution (as already shown on alogarithmic scale in Fig. 2). The same data plottedlogarithmically are shown in Fig. 4 and have nonlinearrelationships. Currents measured at other potentials(E=0, 0.1 and 0.2 V vs. SCE) show a similar pattern.This implies that the apparent reaction order withrespect to hydrogen peroxide, in its oxidation reactionof the first anodic wave, is dependent on its ownconcentration. It does not depend on potential, or inother words, reaction orders are the same in the rising

the aid of ring-disc voltammetry (glassy carbon disc,platinum ring), no other reaction product besides oxy-gen could be determined. Moreover, gas formation atthe glassy carbon surface could be detected visually.Given the composition of the solution (H2O, H2O2 andNaOH), no gas other than oxygen is conceivable as theproduct of an oxidation reaction. The global reaction inthe first wave is therefore supposed to be

HO2−+OH−UO2+H2O+2e− (1)

The transfer coefficient values obtained for the investi-gated combinations of hydrogen peroxide concentra-tions and pH are given in Table 1. All values comeclose to 0.5.

In a second series of experiments, oxidation waves ofhydrogen peroxide were recorded at a rotating glassycarbon disc electrode at a temperature of 298.0 K anda constant rotation speed of 1000 rpm for a relativelylarge number of pH values (10–14 with steps of 0.1pH-unit) and several discrete hydrogen peroxide con-centrations (0.02, 0.05, 0.1, 0.2, 0.3, 0.4, 0.5 mol dm−3).For all hydrogen peroxide concentrations, currents atfixed potentials, as well as the rising part of the wave(e.g. measured at 0, 0.1 and 0.2 V vs. SCE) as in thepseudo limiting region (measured at 0.45 V vs. SCE),are proportional to the concentration of hydroxyl ionsin solution. These ions therefore take part in the elec-trode reaction. This is shown in Fig. 2, where thelogarithm of the pseudo limiting current at a potentialof 0.45 V versus SCE is plotted versus pH. The slope ofthe linear relationships, that reflects the reaction orderof OH−-ions, decreases (from 0.54 to 0.30) with in-creasing (from 0.02 to 0.50 mol dm−3) hydrogen perox-ide concentration. In other words, the reaction order ofOH−-ions depends on the concentration of hydrogenperoxide.

Fig. 3. Dependence of current signal at E=0.45 V versus SCE onhydrogen peroxide concentration for the first oxidation wave ofhydrogen peroxide at a glassy carbon electrode; T=298.0 K andN=900 rpm. pH values are (x) 10.0; ( ) 10.5; () 11.0; (2) 11.5;(") 12.0; (�) 12.5; (�) 13.0; (�) 13.5; () 14.0.

P. Westbroek, E. Temmerman / Journal of Electroanalytical Chemistry 482 (2000) 40–47 43

Fig. 4. Logarithmic plot of the relationship between current signal atE=0.45 V versus SCE and hydrogen peroxide concentration for thefirst oxidation wave of hydrogen peroxide at a glassy carbon elec-trode; T=298.0 K and N=900 rpm. pH values are (x) 10.0; ( )10.5; () 11.0; (2)11.5; (") 12.0; (�) 12.5; (�) 13.0; (�) 13.5; ()14.0.

neously in the anodic prewave in the whole range ofhydrogen peroxide concentrations studied. The varia-tion of reaction orders with concentrations of hydrogenperoxide could then be explained by the followingassumptions:1. two mechanisms are operational2. they are in competition with each other and their

relative contribution in the global reaction dependson the hydrogen peroxide concentration

3. reaction orders with respect to OH−-ions and hy-drogen peroxide in each mechanism lie outside therange of experimentally obtained orders and at op-posite ends of it.

The reason for the requirement that real reactionorders lie outside the experimental range is that theexperimentally obtained orders still vary considerablywith hydrogen peroxide concentration at the outer lim-its of the concentration range of hydrogen peroxidestudied. This means that both postulated mechanismsare operational for all experimental data and the situa-tion whereby the influence of one or both mechanismscan be neglected lies beyond experimentally accessibleconditions.

No direct evidence of the intermediates involved inthe oxidation reaction was found by cyclic voltammet-ric and ring-disc voltammetric experiments. As elabo-rated further, the experimental evidence, Tafel slopeand apparent reaction orders are compatible with thecombination of the pair of rival mechanisms shown inScheme 1.

In step 1 (Eq. (2)) of both mechanisms, hydrogenperoxide is adsorbed at the surface of the glassy carbonelectrode. HO2

−, unstable because of the localized nega-tive charge at the terminal oxygen atom, is stabilized byadsorption. Adsorbed O− particles and OH radicals areformed in steps 3 and 4 (Eqs. (3) and (4)) of bothmechanisms. The resulting degree of surface coverageexplains the experimentally observed small current den-

part of the wave and in the pseudo limiting currentrange. For the lowest used hydrogen peroxide concen-tration the reaction order varies between 0.49 (pH 14)and 1.22 (pH 10); for the highest concentration thatwas studied the order varies in the 1.05 (pH 14)–1.28(pH 10) range.

From the results described above it is clear that acomplicated mechanism is governing the oxidation re-action of hydrogen peroxide at the surface of a glassycarbon electrode in the first anodic wave.

4. Discussion

The puzzling variation of the apparent reaction orderwith respect to OH−-ions with hydrogen peroxide con-centration and of the order with respect to hydrogenperoxide with both hydrogen peroxide and OH− con-centrations led us to the hypothesis that at least twodifferent mechanisms must be operational simulta-

Mechanism BMechanism A

HO2−U

k1

k−1

(HO2−)ads (2) HO2

−Uk1

k−1

(HO2−)ads (2)

(HO2−)adsU

k2

k−2

(O−)ads+(OH·)ads (3) (HO2−)adsU

k2

k−2

(O−)ads+(OH·)ads (3)

(OH·)ads+OH−Uk3

k 3

(O−)ads+H2O (4)(OH·)ads+OH−Uk3

k 3

(O−)ads+H2O (4)

(O−)ads+HO2− U

k4%

k−4%

(O2−)ads+OH− (5%)(O−)ads+(O−)adsU

k4

k−4

(O2−)ads+e− (5)

Scheme 1.

(O2−)adsU

k5

k−5

O2+e− (6) (O2−)adsU

k5

k−5

O2+e− (6)

P. Westbroek, E. Temmerman / Journal of Electroanalytical Chemistry 482 (2000) 40–4744

sities. The fourth steps of both mechanisms (Eqs. (5) and(5%)) are the only steps that are not the same and reflectthe competition between both mechanisms. In mechanismA, adsorbed oxygen anions react with each other, whilein mechanism B they react with hydrogen peroxide anions.This implies that by increasing the hydrogen peroxideconcentration, mechanism B becomes more importantand, hypothetically, it could gain complete control of theglobal reaction at sufficiently high, but experimentally notaccessible, hydrogen peroxide concentrations. On theother side, the situation that only mechanism A isoperational, would occur at presumptive extremely smallhydrogen peroxide concentrations. Recalling that thereaction order, with respect to OH−-ions, is 0.54 at thelowest and 0.3 at the highest hydrogen peroxide concen-tration studied, theses orders in mechanisms A and Bshould meet the following requirements. In mechanismA, which predominates at lower hydrogen peroxideconcentrations, the reaction order with respect to hy-droxyl ions should be higher than 0.54. In mechanism B,prevalent at higher hydrogen peroxide concentrations,this order should be lower than 0.3. An additionalrequirement is the immutable value of 0.5 found for thetransfer coefficient. The theoretical validity of the pro-posed mechanism pair, with respect to these constraints,was verified in the following way. The global rate of eachof the mechanisms is determined by its rate determiningstep (RDS). The rates of the individual steps, when theyare not rate determining are supposed to be in quasiequilibrium. The rate equations of the distinct steps are

61=k1cHO2−(1−uHO2

−)−k−1uHO2− (7)

62=k2uHO2−(1−uO−)(1−uOH·)

−k−2uO−uOH·(1−uHO2−) (8)

63=k3uOH·cOH−(1−uO−)−k−3uO−(1−uOH·) (9)

64=k4(E)uO−2 (1−uO2

−)−k−4(E)uO2−(1−uO−)2 (10)

64%=k4%cHO2−uO−(1−uO2

−)−k−4%cOH−uO2−(1−uO−)

(11)

65=k5(E)uO2−−k−5(E)cO2

(1−uO2−) (12)

where the subscript of 6 indicates the step (Eqs. (2)–(6)),u is the surface coverage and k, a rate constant. A forwardreaction (oxidation for electrochemical reaction) is differ-entiated with respect to a backward one by a minus signbefore the subscript. For electrochemical reaction steps,the rates are potential dependent and their constantwritten as k(E) (to keep equations concise, the explicitpotential dependence is introduced further).

Coverage values u, that are all unknown, can be partlyeliminated from each of the Eqs. (7)–(12) by supposingthat all steps, except the rate determining one, are invirtual equilibrium, i.e. by use of one or more of thefollowing equations that express the condition of quasiequilibrium for the reaction steps considered when they

are not rate determining

k1cHO2−(1−uHO2

−)=k−1uHO2− (13)

k2uHO2−(1−uO−)(1−uOH−)=k−2uO−uOH−(1−uHO2

−)

(14)

k3uOH·cOH−(1−uO−)=k−3uO−(1−uOH·) (15)

k4(E)uO−2 (1−uO2−

)=k−4(E)uO2−(1−uO2

−)2 (16)

k4%cHO2−uO−(1−uO2

−)=k−4%cOH−uO2−(1−uO−) (17)

k5(E)uO2−=k−5(E)cO2

(1−uO2−) (18)

With the aid of the appropriate equilibrium equations,the rate equations (Eqs. (7)–(12)) can be transformed intoEqs. (19)–(23) and Eqs. (24)–(28), respectively for mech-anism A and B.

61,A=k1cHO2−−

1K2K3

k−1k %−4k %−5

k %4k %5

cO2

cOH−

e−2F

RT(E−Eref)

(19)62,A=K1k2cHO2

−(1−uHO2−)

−1

K3

k−2k %−4k %−5

k %4k %5

cO2

cOH−

e−2F

RT(E−Eref) (20)

63,A=K1K2

k3k %41/2k %51/2

k %−41/2k %−5

1/2

cHO−

cO21/2 (1−uOH·)e

F

RT(E−Eref)

−k−3k %−4

1/2k %−51/2

k %41/2k %51/2 cO2

1/2eF

RT(E−Eref) (21)

64,A=K1K2K3k %4cHO2−cOH−(1−uO−)2e0.5

F

RT(E−Eref)

−k %−4k %−5

k %5cO2

e−1.5F

RT(E−Eref) (22)

65,A=K1K2K3

k %4k %5k %−4

cHO2−cOH−(1−uO2

−)e1.5F

RT(E−Eref)

−k %−5cO2e−0.5

F

RT(E−Eref) (23)

61,B=k1cHO2−−

1K2K3K %4%2

k−1k %−52

k %52cOH−cO2

2

cHO2−

2

�

e−F

RT(E−Eref)�2

(24)

62,B=K1k2cHO2−(1−uHO2

−)

−1

K3K %4%2k−2k %−5

2

k %52cO2

2 cOH−

cHO2−

2

�e−

F

RT(E−Eref)�2

(25)

63,B=K1K2K %4%k3k %5k %−5

cHO2−

2

cO2

(1−uOH·)eF

RT(E−Eref)

−1

K %4%

k−3k %−5

k %52cO2

cOH−

cHO2−

eF

RT(E−Eref) (26)

64%,B=K1K2K3k %4%cHO2−

2 cOH−(1−uO−)

−k %−4k %−5

k %5cO2

cOH−e−F

RT(E−Eref) (27)

P. Westbroek, E. Temmerman / Journal of Electroanalytical Chemistry 482 (2000) 40–47 45

65,B=K11/2K2

1/2K31/2K %4%k %5

cHO2−

3/2

cOH−1/2 (1−uO2

−)

×e0.5F

RT(E−Eref)−k %−5cO2

e−0.5F

RT(E−Eref) (28)

where

kx(E)=k %xe(1−a)F/RT(E−Eref)

and

k−x(E)=k %−xe−aF/RT(E−Eref)

Each of these equations is composed of two terms. Thefirst term relates to the forward reaction or the oxidationof hydrogen peroxide, the second, negative term, corre-sponds to the backward reaction. In the reaction studied,the cathodic part can be neglected completely because ofthe irreversibility of the HO2

−/O2 redox system and therelatively positive applied potentials. Looking moreclosely at the anodic terms, it is clear that the possibilityof steps 1 and 2 of mechanism A and steps 1, 2 and 4of mechanism B acting as the RDS can be eliminatedbecause their reaction rates are independent of potential.This is in contradiction with the results obtained exper-imentally. Also step 3 of both mechanisms can beeliminated as a possible RDS because the reaction rateis dependent on oxygen concentration, which has not beenobserved in the experiments. Using the transfer coefficientof 0.5 as a criterion in the rising part of the oxidationwave, the fifth step of mechanism A can also be eliminated.

Remaining as a possible RDS are step 4 (Eq. (5)) ofmechanism A and step 5 (Eq. (6)) of mechanism B, thatcomply with the requirements expressed previously. Thepredicted transfer coefficient of 0.5 is in agreement withthe experimental value in the rising part of the wave. Step4 of mechanism A and step 5 of mechanism B predictreaction orders of, respectively 1 and −0.5, with respectto OH−-ions, in agreement with the above deducedrequirement of being higher than 0.54, and lower than0.3, respectively. The combination of mechanisms A andB predicts that the reaction order of hydrogen peroxidewill increase its concentration from 1 for mechanism A,at very low concentrations, to 1.5 for mechanism B at highconcentrations.

The formation of a pseudo limiting current plateau canbe explained by the fact that in mechanism A, 63,A (Eq.(21)) is inversely proportional to the square root of theconcentration of dissolved oxygen and 64,A (Eq. (22))varies linearly with (1−uO−) and that in mechanism B,63,B (Eq. (26)) is inversely proportional to the concentra-tion of dissolved oxygen and 65,B (Eq. (28)) varies linearlywith (1−uO2

−). The oxygen two phase system (dissolvedoxygen/oxygen in the gas phase) at the solution � electrodesurface interface appears to play a key role in theformation of a pseudo limiting current plateau.

Eq. (21) shows that an increase of 63,A with potential,with an exponent twice as large as that of 64,A (Eq. (22))

is attenuated by the presence at the electrode surface ofdissolved oxygen. The latter is the product of the globaloxidation reaction (Eq. (1)), so more of it is present atmore positive potentials and higher bulk concentrationsof hydrogen peroxide. When the solubility of dissolvedoxygen at the electrode surface is exceeded, the rate offormation of O− in step 3 (Eq. (4)) increases more rapidlythan the rate of its conversion into O2

− in step 4 (Eq. (5)).O− piles up at the electrode surface and (1−uO−) in Eq.(22) decreases and for the most part, counterbalances theexponential increase of its rate with potential. Thishypothesis of the origin of a kinetically controlled pseudolimiting current is supported by the observation that theslope of the pseudo limiting current plateau is higher(counterbalance of both effects is not complete) and Tafelbehavior extends to more positive potentials when thehydrogen peroxide concentration, and thus the concentra-tion of dissolved oxygen, at the electrode surface arelower. For hydrogen peroxide concentrations exceeding2 mol dm−3, the plateau slope becomes negative, indicat-ing an overbalance of the exponential current increase.

A largely similar reasoning can be made for mechanismB. Here the exponential increase with the potential of ratedetermining step 5 (Eq. (28)) is to a large extent neutral-ized by the decrease of (1−uO2

−), again induced by theeffect of the transgression of oxygen solubility on 63,B (Eq.(26)) and secondarily on 64%,B.

In a following step in the verification of the validityof the emitted hypothesis, the reaction rates of thecompeting mechanisms have to be unified into a singleglobal rate, because experimentally only one, global,current can be measured. This is possible by introducinga parameter x that marks the relative share of mechanismA in the total current and depends only on the concen-tration of hydrogen peroxide in solution. Parameter xwould equal one if mechanism A (with step 4 as the RDS)were to be operational only or, in other words, atextremely small hydrogen peroxide concentrations, wherea reaction order of one with respect to both hydrogenperoxide and hydroxide is expected (Eq. (22)). Reversiblyx would equal zero if mechanism B (with step 5 as theRDS) would be operative only, corresponding to thesituation of extremely high hydrogen peroxide concentra-tions and a reaction order of 1.5 with respect to hydrogenperoxide and −0.5 with respect to OH−-ions (Eq. (28)).With the aid of the parameter x, Eqs. (22) and (28) canbe combined into

I=G(x, E)cHO2−

(1.5−0.5x)cOH−(1.5x−0.5) (29)

where

G(x, E)=nFAK(1−uO−)2(1−uO2−)e0.5

F

RT(E−Eref) (30)

and

K=K1(0.5x+0.5)K2

(0.5+0.5)K3(0.5x+0.5)K %4(1−x)k %4xk %5(1−x)

(31)

P. Westbroek, E. Temmerman / Journal of Electroanalytical Chemistry 482 (2000) 40–4746

Fig. 5. Relationship between parameter x (Eq. (32)) and hydrogenperoxide concentration for the oxidation of hydrogen peroxide at aglassy carbon electrode at E=0.45 V versus SCE and T=298.0 K.

Fig. 6. Comparison of hydrogen peroxide concentrations obtained bycontinuous measurement of the amperometric current at a glassycarbon electrode at E=0.45 V versus SCE (continuous line) withtitrimetric results (x, , �).

where Kn=kn/k−n. The value of G(x, E) at a potentialof 0.45 V is designated by G %(x).

In order to verify if experimental currents obey Eq.(29), the relationship between x and the hydrogenperoxide concentration should be known. From theexperimentally obtained apparent reaction orders andthe theoretically predicted ones (equal to 1.5x−0.5, seeEq. (29)) with respect to hydroxyl ions (which aredependent on the hydrogen peroxide concentration, Fig.2), x values were determined for a series of hydrogenperoxide concentrations. The results are shown in Fig. 5.From a logarithmic plot of the data one derives

x=0.5071cHO2−

−0.0791 (32)

In the hydrogen peroxide concentration range studied(0.02–0.50 mol dm−3), the value of x varies from 0.691to 0.535.

Once the value of x is known for a series of hydrogenperoxide concentrations, G(x, E) at any selected poten-tials can be calculated from current data, pH beingknown, with the aid of the logarithmic version of Eq. (29)

log I= log G(x, E)+ (3/2−1/2x)log cHO2−

+ (3/2x−1/2)log cOH− (33)

G values obtained in this way at a potential of 0.45 Vversus SCE for different hydrogen peroxide concen-trations, lead to the following relationship between G %and x

G %=G(x, 0.45 V vs. SCE)=9.3×105x6.607 (34)

Eqs. (29)–(31) allow us to understand qualitatively theexperimental relationships shown in Figs. 2–4. As bothG(x) and x depend on hydrogen peroxide concentration,

linear and logarithmic calibration plots are curved (Figs.3 and 4) at all pH values. The pronounced influence ofthe pH itself and the fact that its impact depends onhydrogen peroxide concentration, follow directly fromEq. (29) as both G(x) and x depend on this concentra-tion.

Finally the influence of temperature was investigatedand it was found that the signal is exponentially depen-dent on temperature according to

I=k %e0.006422T (35)

where k % is a constant value.The validity of Eq. (29) was verified by simulating its

use for analytical purposes by determining hydrogenperoxide in a wide range of and up to large (about 1 moldm−3) concentrations. The glassy carbon electrode wascalibrated by measuring the amperometric current at aknown pH and temperature and determining the hydro-gen peroxide concentration by titration with potassiumpermanganate in acidic medium. Then the amperometriccurrent was measured in a solution that containedhydrogen peroxide in a concentration supposed to beunknown. This concentration was also determined titri-metrically, in order to compare titrimetric and ampero-metric results.

When Eq. (29) is written twice, once for a solution usedin calibration and once for a solution with an unknownhydrogen peroxide concentration, the latter can be ex-tracted from both equations as

where the subscripts cal and m refer to the solution usedfor calibration and the solution in which the hydrogenperoxide concentration has to be determined, respec-tively, p=0.5071 and q= −0.0791. Hydroxyl ion con-

cHO2,m− =

�cHO2,cal

−(1.5−0.5pcHO2,cal

−q ) Im

Ical

cOHcal−

(1.5pcHO2,cal−q −0.5)

cOHcal−

(1.5pcHO2,m−q −0.5)

(pcHO2,cal−

q )6.609

(pcHO2,m−

q )6.609 e0.00642(Tcal−Tm)n1/(1.5−0.5pcHO2,m−

q )

(36)

P. Westbroek, E. Temmerman / Journal of Electroanalytical Chemistry 482 (2000) 40–47 47

centrations are obtained from pH measurements inboth solutions. The unknown hydrogen peroxide con-centration was obtained by iterations using Eq. (36).Some of the results obtained by continuous measure-ment of the amperometric current and periodical titra-tion of an aliquot of the solution are summarized inFig. 6. The hydrogen peroxide concentrations obtainedby amperometry deviate by less than 2% from resultsobtained by titration.

5. Conclusions

A possible mechanism has been derived for the oxi-dation reaction occurring in the first oxidation wave ofhydrogen peroxide at a glassy carbon electrode in alka-line solutions. This mechanism is composed of twosubmechanisms, which are in competition with eachother. This competition can explain why the reactionorder, with respect to hydroxyl ions is dependent on thehydrogen peroxide concentration. The occurrence ofadsorbed intermediate species can explain why apseudo limiting current is observed and why the currentdensity is relatively small. From the current equations,for each mechanism an equation was derived for thecalculation of unknown hydrogen peroxide concentra-tions from measurements of amperometric current sig-nal, pH and temperature.

The results described in this paper open possibilitiesfor the use of the reaction studied in a sensor system fordirect and continuous monitoring and for the control ofconcentrations of hydrogen peroxide up to high values.

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