Chronopotentiometric reduction of oxygen at platinum electrodes in sulfuric acid media

JOGRSXL

306

CHRONOPOTENTIOMETRIC ELECTRODES IX SULFURIC

DEKXIS

G_ PETERS

23epatiment (Received

_4?iD

ROXXLD

of CIzemisfq~. Indiana April

REDUCTION ACID MEDIA*

A.

Grzruusity,

OF

ELECTROASILE-TICAL

OF

O_XE’GEN

CHEMISTRl-

_AT

PLATINUM

MITCHELL Bloomington,

Inrlm~~a

(U.S._4

_)

6th. 1965)

The electrochemical reduction of dissolved oxygen at platinum electrodes has of the reduction as the subject of a number of recent studies i--6_ The mechanism well a_s the influence of electrode surface condition on the course of the reaction has been of interest. The present report is devoted to chronopotentiometric and chemical investigations of o_uE_Sen reduction at platinum electrodes in sulfuric acid media and to the inter-relationship between electrode surface condition and oxygen reduction_ Because the surface condition of a platinum electrode does influence the process of oxygen reduction, a brief review of the possible surface states of a platinum electrode seems warranted_ Indirectly, the existence of several surface states of a platinum electrode has been inferred and, at least in part, demonstrated. AXSON ASD KINGS have classified these surface states and described techniques for controlling or altering the surface condition of an electrode_ Chemical oxidation or anodic polarization of a platinum electrode in non-complexing media, e.g., sulfuric acid solutions, produces occur platinum oxide films (oxi&ked platinum surface) s--11_ When electrode reactions at oxidized electrodes, overpotential and/or inhibition effects are encounteredl’-14. Chemical or electrochemical reduction of an oxidized platinum electrode results in the formation of a deposit of finely-divided, catalytically-active platinum metal been

~7es?zZy re&tced or active platinum surface) _ Electrode reactions proceed with the greatest degree of reversibility at freshly-reduced or active electrodes_ Several studies have demonstrated that, upon standing without reactivation, a freshly-reduced or active platinum electrode suffers loss of its deposit of finely-divided metal with resultant formation of an age& vedrcced platinum surface 15_ This aging of the platinum surface reveals itself by increase in the overpotentials for electrode reactions. In the extreme case, a fully aged electrode behaves as the stripped electrode mentioned by ANSON AND KING’_ The chronopotentiometric reduction of oxygen at a fresMy-redzrced or actzve platinum cathode yields a single wave in sulfuric acid media corresponding to the diffusion-controlled, four-electron reduction to water. This conclusion, reached originally -by LINGAXE xX has been corroborated in-the present study. In contrast to the report- by LINGANE, however, we found that, if the active electrode is allowed to stand in the oxygen-saturated solution, it rapidly undergoes an aging process which causes the over-potential for oxygen reduction to increase and the apparent It-value * contribution

In_

(u.s.a_).

No.

1307

from the Department

/_ EZe~~oanaZ- Chem., Io-(;g65)

306318

of Chemistry.

Indiana

University,

Bloom-mgton,

REDUCTIOX

OF

Oz AT Pt-ELECTRODES

Is

HzSO~-MEDIA

307

for the reduction to decrease to tu-o, i-e_, reduction of oxygen only to hydrogen peroxide. On the basis of chronopotentiometric and voltammetric data, SAWYER and collaborators”*3 concluded that the reduction of oxygen proceeds by two different mechanisms_ At @e-oxidi_zecZ platmum electrodes, the reduction of oxygen was postulated to occur by a cyclic process which involves four electrons per mole of oxygen (reduction to water) along I\-ith the direct chemical reaction of oxygen with platinum. For pre-redzlced electrodes, a second mechanism was proposed in which osygen is reduced in a two-electron step to hydrogen peroxide. It should be pointed out that the pre-oLdized electrodes used by SXWYER and co-workers would have become actiae during and immediately following a cathodic trial; the pre-reclrrced cathodes employed by these workers presumably correspond to the aged, renzrcen electrodes described by AXSOS XXD KISG7. The chronopotentiometric and chemical experiments described below suggest that a single mechanism for oxygen reduction is operative for both nctiae and aged electrodes_ We postulate that osygen is electrochemically reduced to h\.drogen peroxide in a diffusion-controlled, two-electron step_ The fate of the hydrogen peroxide is deter-m!& ed by the surface condition of the platinum cathode. _%t a fres?rZ_~ redztced or nctize electrode, hydrogen pero.xide undergoes a surface-catal_vzed disproportionation, as rapidly- as it is generated, to water and more osygcn ; therefore, the pz-value for reduction of oxygen appears to be four. As the electrode ages, however, the ability of the electrode to catalyze the disproportionation of hydrogen peroxide is greatly diminished until finally the reaction becomes so slow that the apparent Tt-value becomes two. EXPERIMENTAL

All solutions were prepared from triply-distilled water, the second distillation being from alkaline perrnanganate, and analytical-reagent-grade sulfuric acid. High purity (gg_6%) oxygen gas was used throughout these experiments; the major impurities were nitrogen and argon. In the few experiments in which oxygen-free solutions were required, pre-purified (gg_gSy/o) nitrogen was employed for deaeration. Titanium(W) sulfate solutions were prepared by appropriate dilution of the solution obtained by dissolution and decomposition of twice-recrystallized, technicalgrade K2TiO(CzO& - 2Hz.O in concentrated sulfuric acid. Conventional chronopotentiometric experiments were performed using an electrolysis cell and electrical circuit similar to those described by LING_~NE~~. Micro working-electrodes were fabricated from lengths of pure (gg_gg”/o) platinum wire of o_ogr cm diameter by coating them with Tygon pamt (Type K-63, U. S. Stoneware Co.. &on, Ohio) in a manner which left exposed a cylindrical area of 0.1-0.3 cma close to one end of each wire. The exact geometric area of each electrode was determined with the aid of micrometer calipers and a binocular microscope. For special experiments conducted to detect and determine hydrogen peroxide formed by reduction of oxygen, several modifications were made in the simple chronopotentiometric apparatus: a large platinum-gauze electrode (approximately IOO cm2 in area) served as the working cathode; a Sargent Coulometric Current Source, Model IV (E_ H. Sargent and Co., Chicago, Illinois), was substituted for the usual battery J_ Electvoesral.

CAem..

IO (1965) 306-318

308

D.

G.

PETERS,

R.

A.

JIITCHELL

current source;

and an annular cell /Fig. I), machined from Teflon to accommodate the gauze electrode, was used in place of the con\-entional chronopotentiometric cell to provide a large electrode-area-to-solution-volume ratio favorable for the subsequent spectrophotometric analysis of dilute hydrogen pero_xide solutions-

Plolrnum

gouze

elecrrode

AlI chronopotentio~ams were recorded on a Varian F-So X-Y recorder (Varian Associates, P&lo Also. California), Unless specificalIy stated otherwise, it may e xperimental data pertain to sulfuric acid solutions through which be assume&&at oxygen wasbnbbled Ior at least 30 min prior to the recording of the chronopotentiogram, A& is__custo_mary, before the chronopotentiogram was recorded the stream of -oxygen was diverted over the surface of the solution, and the test solution itself waS -allowed on& minute or more to become quiescent_ Experiments carried out with +?~e ~&-+&&I electiolysis -Cell were performed at z~.o-&o.I~. Because of the. in~&GGznieii6e of th&nostatting the annular ceh, work with the latter was-done at the pr&+ng_labdr&tory tetipera_ture_ .-_ . : - -% _ y-EZei%op&Z-%he+., x6(1&) 3c$-318

OF OS AT

REDUCTIOS RESULTS

ASD

platinum a current

IS

H&04-MEDIA

309

DISCUSSIOX

Clzro1lopote?ltiolrzetric X

Pt-ELECTRODES

determination

of n-value

for

oxygen

redzrcfiom

pair of cathodic chronopotentiograms for the reduction of oxygen at a electrode in I F sulfuric acid is shown in Fig. 2. Each curve was recorded for density of 6go PA/cm” and, from data given by SEIDELL~~, the concentration

of dissolved

oxygen

in I F sulfuric

acid at one atmosphere

pressure

was calculated

to

TIME

Fig. 1. Chronopotentiograms for osygcn reduction at a platinum-wire cathode in oxygen-saturated I F H2SOI; current density = 6go LrX/crn’. (I), Ox>-gen reduction at an oxidized electrode; prior to the trial, the electrode w-as anodized for 10 set at IO mX/cm’; (2). oxygen reduction at an active electrode, recorded immediately following curve I.

be 1.06 x 10-3 &I. The characteristics of such chronopotentiograms have been discussCurve I of Fig. z is the chronopotentiogram for oxygen ed in detail by LIKGAXE~. reduction obtained with an electrode which was anodized for 20 set at a current density of IO mA/cm” just prior to the trial. The shallow minimum in this curve following commencement of the electrolysis is due to the overpotential associated with oxygen reduction at an oxidized electrode. However, once part of the platinum oxide film has been reduced to active platinum metal, oxygen reduction occurs more reversibly and the potential of the electrode becomes slightly more anodic. Curve 2, to the reduction of oxygen at a recorded immediately after curve I, corresponds freshly-reduced or active platinum electrode_ The enhancement of the cathodic transition time due to the concomitant reduction of the platinum-oxide film is evident from a comparison of curve I with curve 2. The small, poorly-defmed post-waves wave are attributable to the formation of following the main oxygen-reduction adsorbed hydrogen on the electrode surface. LIXGXWE~ demonstrated that the transition time for oxygen reduction at a freshly redwed or active platinum electrode (curve z, Fig. 2) in I F sulfuric acid corresponds to the diffusion-controlled, four-electron reduction to water. This result was confirmed in the present study with transition-time data obtained according to the procedure used by LIXGANE (cf_, ref. I, p_ 303). We performed similar measurements for oxygen reduction at active electrodes in 4 F sulfuric acid in which the solubility of oxygen is o-69 x 10-3 M le_ The diffusion coefficient of oxygen in 4 F sulfuric acid _T_Eleclroaxal.

Ghem = IO (1965)

306-318

D. G. PETERS, R_ A_ MITCHELL

310 was evaluated by conventional polarographic techniques t=3.65 set) ancl calculated from the Lingane-Loveridge cmz/sec at 25O. This value is approximately one-half

(ia = 12_3,uA, 922= 3-05 mg/sec, equation17 to be 1.05 x IO-J of that (2-02 x 10-5 cm’/sec)

determined for I F sulfuric acid b>- LINGXXE~_ From chronopotentiometric 4 F sulfuric acid, we conclude that oxygen undergoes a four-electron, controlled reduction at an act&e electrode.

data for diffusion-

More detailed studies of oxygen reduction revealed a strong inter-dependence between the character of the wave and the electrode surface condition. In particular, as the elapsed time from the original activation (or reduction of the platinum oxide film) increases, the overpotential for oxygen reduction increases and, simultaneousl)-. the transition time for the cathodic wave decreases. However, the oxygen-reduction wave does not completely disappear as LIXG_&WE~ reported, but diminishes only until the transition time corresponds to the diffusion-controlled, two-electron reduction of oxygen

to hydrogen

peroxide.

TIME

Fig. 3_ Effect

of aging of a platinum electrode on osygen reduction in 4 F H&O,The elcctmde activated (anodized and cathodized) and allowed to stand in the o.xygen-saturated solution. WaS _4t various times following activation, a chronopotentiogram was recorded at a current density of 310 PA/cm”. (a), z min; (b). 15 min; (c), I h 15 min; (d). 5 h; (e), 24 h; (f), 127 h after activation.

Figure 3 shows this sequence of events oxygen in 4 Fsdfuric acid. Prior to the recording

for the reduction of the first curve

of o-69 x 10-3 _ii in Fig. 3, the plati-

num-wire electrode-was anodized (20 set at 2.2 mA/cmz) and then cathodized just to the hydrogen evolution potential at the same current density to remove the oxide film. This procedure produces what we have referred to previously as an actzve electrode, because the over-potential for oxygen reduction is rnintial and because, as shown below, hydrogen peroxide c&proportionates rapidly on the electrode surface resulting from this pretreatment_ After this activation the electrode was allowed to stand in the sulfuric acid solution ; then, at various intervals following this activation, cathodic chronopoten~ograms for oxygen reduction were recorded. The apparent rt-value for each oxygen cbronopotentiog-ram obtained during this agingiequb+ was caIculated from the PetersiLingane equationI*; the measurement of the transition times-demanded special attention because of the shift in the potential of -the_ oxygen wave as the electrode ages. Therefore, the transition times were evaluated from the complete pen-and-ink recording of the chronopotentiograms (Fig. 3) - _ J_ &ectma~zaZ.

.&em..-ro

(1965)

3ck318

REDUCTIOK

0~;

02 AT I?-ELECTRODES

Ih’ H2S04-~rE~~a

3==

by measurement of the time elapsed from the beginning of the electrolysis to a potential selected as the point of inflection of the individual chronopotentiogram_ (For the curves in Fig. 3, these potentials are, in order, +O.IO, + 0.10, + 0.10, o, -0.06, and -O.IZ

potential

V “JS_S.C.E.) Measurements of the transition times at the same preselected of + 0.10 V 2s. S.C.E. apparently led LIKGAKE~ to conclude that the ?z-value

decreased

all the way

In connection

to zero. with

the series

of chronopotentiograms

shown

in Fig.

3. several

additional remarks should be made. First, the six chronopotentiograms of Fig. 3 were selected from a total of nineteen recorded curves in one particular experiment in order to depict the important consequences of the aging process_ Second, the aging of the electrode, or the decrease in the n-value, depends almost entirely on the time elapsed from the original activation of the electrode and is largely independent of the number of cathodic chronopotentiograms recorded. However, when a chronopoten tiogram is recorded, the electrode is re-activated to a slight extent, especially if the electrode has reached a \-ery aged condition as for the last two curves in Fig. 3; that is, the overporential for oxygen reduction is somewhat decreased and the transition time is observed this same kind of catkodic re-activation and slightly increased. LIKGXXE1 attributed it to the reduction of platinum ions (from dissolution of a chemicallyformed platinum oxide film) and the resultant renewal of the deposit of finely-divided We found that the extent of cathodic platinum metal on the electrode surface. re-activation increases as the current density and duration of polarization increase; therefore, catkodzc re-activation could be associated with formation of adsorbed atomic hy-drogenlg.

I

I

I

I

I

I 40

I

I

I

I

I

60

80

100

120

140

I

I

I

I-

0

20

ELAFSED

Fig. time (O),

+ Effect elapsed 4 F

of aging of a platinum from original activation

TIME,

HOURS

electrode on oxygen of an electrode in

reduction. F and

I

4

Plot of apparent F H&O,_ ( A),

I

x-value x. F H&04;

H.&h-

Figure 4 shows a plot of the apparent x-values for successive cathodic chronopotentiograms versus time elapsed from the original activation of the electrode for oxygen-saturated I F and 4 F sulfuric acid media. The rapid drop of the apparent J. EZcctroanaZ. Chern..

IO (1965)

30631s

3x2

D.

C_

PETERS, R_ -1. XITCHELI.

S--c-aluc from four to approximately two, whence the electrode then ages orders of ~;lower, is consistent \%-ith the hypothesij: that, at a freshly-acti!:ntecl platinum surf ace, the osygen is reduced to hy-dmgen peroxide which then dispro215 the electrode age”, the platinum surface portionates to water and more oxygen. loses this ability to catalyze the disproI)nrtirtnation of Ii>-drogcn perosicle and the reduction of osygen proceeds on:y as far as hydrogen peroside. We would emphasize that the results sho\\n in Figs. 3 and 4 are typical and that the overall trends are readily reproducible_ Nevertheless, although the influence of the electrode aging process on oxyg en reduction appears to be established, quantitative differences in the rate of aging are observed under presumably identical experimental conditions. Obviously, the factors which govern the rate of aging are not fully understood. Further research is being directed toward this goaI. The fact that an activated ekctrode does have the ability to catalyze the di+ propo~iona~o~ of hydrogen peroxide, but that the electrode loses this ability upon aging, was demonstrated in the following large-scale experiments. Details of the procedures used to activate and age the platinum gauze electrode are described below. An aliquot of 0.171.F hydrogen peroxide in I: F sulfuric acid was pipetted into the annular cell (Fig I) which already contained the j&sMy a&vnterZ platinum-gauze electrode. %Vithin seconds, bubbles of oxygen were observed on the electrode surface; two hours later, titrimetric knalysis of the solution with Ce(IV) sulfate showed that 96:/o of the hydrogen peroxide had disproportionatedThe experiment was repeated with the platinum-gauze electrode in an agerl condition similar to that of the platinum-wire electrode of the last two cumes of Fig. 3. No bubbles ok oxygen were seen on the efectrade and, when the solution was analyzed after being in contact with the electrode for two hours, hydrogen peroside Was found to be only 3yG disproportionated, Areportby LINGANEAXD LINGAXE'O on the chronopotentiometry of hydrogen peroxide in sulfuric acid medium provides additional evidence that an nctk~e platinum surface Catalyzes the rapid disproportionation of hydrogen peroxide while an n,ozrE electrode does not. As an origiirally active electrode ages and loses the ability to catalyze hydrogen peroxide disproportiontition, it might be expected that a separate wave for reduction .of- hydrogen, peroxide to water wouId be observed_ In other words, oxygen should undergo stepwisc reduction, first -to hydrogen peroxide and then to water, at an aged ele&o&. However, no clear indication of a wave for the reduction of hydrogen peroxide’to virater’was qbtained. The explanation for this behavior appears to be that for an cagedelectrode the overpotent& fdr reduction of. hydrogen peroxide to water is so large th& any second wave is masked by hydrogen ion reduction Indeed, when a pfatintim~wire electrode was aged in oxygen-saturated4 Fsutfuric acid until the +value ‘reached t&o and a‘ten:fold-excess of hydrogen.peroxide waS added to the solution, the subkquent c~th~~~c~n~~ot~~~o~ kas vir.tual.ly identical to thatobtained in the .tibsence df hydrogen peroxide..+ the study of thechronopotentiometric reduction of hydrogenperoxide; ~.~~GANE.$ND LIXck~tiReobt&ned c~ono~ot~n~io~~s usingpar&zZ@ q&d, platinurne~ectrodks wbkh exhibited &&rate reduction waves for oxygen (G+ich cj+ini$esfr$in the. su.rfac&atalyzed disproportion&ion“ of hy_drdgen peroxide) and for’hydrogen. pe&kide; ,I&’ the pre_&nt investigation, however, oxygen reduction at’ ‘@g&iiL~ &t+d -eIec~odes (cfr,..‘the:‘second and third- Curves in Fig- 3) did. not produce kIffiki&ntJyY w&l&veloped ~~vaves:to~&ZiCate ‘unequivocally stepwjse reduction; _: magnitude

-~I-Ele’.“o!rsr.:~~~~~;.

IO ‘(f&j,

.-.

3c+G318.-’

1~

HsSOa-3IEDI.A

313

enidelrce for ?t)'cZ~oge~~.peroxide formation at aged electrodes To demorrstrate independently that hydrogen peroxide is a stable product of osggen reduction at ngecl electrodes, chemical (spectrophotometric) analysis of the test solutions was performed_ Single chronopotentiometric electrolyses were carried out in a speciall_v designed annular cell %vhich accommodated a large platinum-gauze electrode (Fig. I). After the chronopotentiometric trial, Ti(IV) was added to the cell solution, -aud the quantity of hydrogen peroxide formed from the reduction of oxygen was determined from the absorbance of the Ti(IV)-peroxide complex at 4ro m,u. The detailed procedure for the determination of the quantity of hydrogen peroxide formed from oxygen reduction is as follows. A\‘ith the platinum-gauze electrode of the desired condition (active or aged) in the annular cell, a cathodic chronopotentiogram was recorded for osygen-saturated I I; or 4 F sulfuric acid. The electrode was immediately removed from the cell and the solution was quantitatively transferred to a 5o-ml volumetric flask. Esactly one milliliter of 0.1 A4 Ti(IV)-4 F sulfuric ac’d solution was added to the flask and the solution was diluted to the mark with I F or 4 F sulfuric acid. The spectrum of the resulting solution was obtained with a Gary 14 recordin g spectrophotometer. Ten-centimeter quartz absorption cells were used- the reference solution contained the same concentration of Ti(IV) and sulfuric acid as the sample sdlution. The concentrations of Ti(IV) in the sample and reference solutions were carefully matched because Ti(IV) begins to absorb at approximately 370 rnp_ From the absorbance measurement at 410 rnp (where the molar absorptivity of the Ti(IV) -peroxide complex is 730 l/mole - cm in both I F and 4 F sulfuric acid) and the quantity of electricity involved in the reduction of osygen, the percentage yield of hydrogen peroxide was calculated. One set of experiments was performed for I F sulfuric acid medium. The large electrode W-U activated (anodization at a current of rg3 mA for zo set followed by cathodizaticu at the same current just to hydrogen evolution), washed with I F sulfuric acid’?..and transferred to the annular cell which contained oxygen-saturated I F sulfuric acid. After the solution was allowed to become quiescent, a chronopotentiogram was recorded in the usual manner (i = 4s mA) . The solution in the annular cell was analyzed for hydrogen peroxide by the procedure described above. A spectrum obtained in a typical experiment is curve 2; of Fig. 5. The bottom spectrum (curve I) of Fig. 5 is the base line obtained with the reference solution in both the ._ sample and reference absorption cells Siuce the amount of hydrogen peroside to be uererrninea was extremely smair, tne recorcung spectropnotometer was equippea wrtn a sensitive slide-wire (o-0.1 optical density unit, full scale) ; therefore, the noise level of the spectra was extremely high. The absorbance at 4ro rn,z in the experiment with the active electrode (curve 2, Fig. 5) corresponds to no more than a 17; yield of hydrogen pero-tide_ Norkover. the absence of a definite peak at 410 my suggests that the small amount of absorption is due to background rather than to the Ti(IV)-peroxide complex _ The sam& experiment was repeated, except that the platinum-gauze electrode was in an aged. condition_ Although a platinum-wire microelectrode ages (*t=a) within 20 h on standing in I F sulfuric acid (Fig. 4), the large gauze electrode displayed little tendency to agein a period of ten days.. The only apparent difference between the w-ire electrode and the gauze electrode during the aging process was the ratio of the,electrode _ zirea to the solution volume. ‘In an effort to provide a large ratio of

Chemical

J- Elect-roanaZ.- Cherw.. .IO (1965) ,306-3x8

314

D.

G.

PETERS,

R.

A.

XITCHELL.

solution volume to electrode area, and hence hasten the aging of the large gauze electrode, the I F sulfuric acid solution in contact with the electrode was changed four tin-r&seach day for ten days. This e,uFeclient was found to age the large electrode quite effectively. Curve ‘3 of Fig. 3 is the spectrum obtained from the experiment with the qe~2 electrode. The definite peak at 410 mp, corresponding to a 33 y/oyield of hydrogen -peroxide, shows that hJrdrogsn peroxide is a stable reduction product of osygen at an a&d. electrode.

303

375

525

450

WAVEIEUGTH,

600

mp

Fig. 5_ Spectro@hotometric evidence for hydrogen peroside formaticn in oqgen reduction at an ngedplatinum electrode in I F N&O+ (I), Base line; (2). spectrum obtained on addition of Ti(IV)

to cathol>*c following of Ti(IV) to catbolyte

I

oxygen reduction at an acLiue electrode; (3). spectrum following o_xygen reduction at an aj& electrode.

obtained

oc addition

The experiments were’ repeated with both active and aged electrodes in 4 F sulfuric’&id, and the. following results were obt.a.ined. With an actiw electrode, little & tie hydrogen peroxide was Petected in the solution. With an aged electrode, yields of hydrogeti-peroxide of 44 , $8, 60, 69, and gI.% were obtained in five different experin..&.+_ ,: : We incliiie to the conclusion that the yields of hydrogen peroxide in all experwith f&z a&d platinum-gauze electrode should have been roo”/& in agreement tith the obsetiation that. the n-value -for .?xygen reduction at an aged pIatinuin rnicrdele~tjFode.~~two_ However, due to the difficuIti&s of ensuring that the large elec. trade. WC& completely $.@ed, thk ‘theoretical yield of hydrogen, peroxide was not obtained.-very likely,,-& significant fractipn of the hydrogen peroxide disproportionated ‘at the inclb~pl&t&Iy a&d electrode.. : About t&enty&ve ye&$ a&o, LAITI~N AT$D KoYkZoFF2l studied the electroehemi&l ~educti& : 0-f &c&en at a. stationary platinum mkroelectr6de in various rried.&using the.volta.mmetr;c technique, They repprted that. hydrogen perotide’is the reduction $rodu& and that. the &enti;il~ at ~~h.i& oxygeti is reduced is independent & hv:.~ v&y .hkely. the :plat++, micrdelectrod&s,‘used. by L;+ITINEN AhTD I~QLT~~OFF w~re_in:;an.~~:cb~~~onj.-aTicl indeed there is,-no. &@i&$orj .th& these.authors at-.. $&t$f~~ .t~[a&iv& :lhe -ek$rodes. ‘Recdn~ly, S+ti-&~k ~,D..I~~~R+LANITP> on- repeat-. i$ig .fhe :&or% ef :~Ii&rr~kti’~&$ Ko&&FF; ~~3jtai+d the same resm&~ p;;ith &a&mm :

imt$iti

REDUCTIO-~;

OF

microelectrodes from

these

OS AT I%-ELECTRODES known

Evidence

of

to be in an aged

voltammetric

chronopotentiometric reduction product

IN HzSO~-MEDI_*

studies

are

(pre-recl~cerZ) state.

compatible

investigation, for we have (?t=z) at an ngsd electrode_

?zy&ogew

$eroxide

as zmstable

3i5

with shown

iaternzediate

The

those that

SC

conclusions

based

on

hydrogen

oxygen

reached

the

present

peroxide

reci!mdio~L

at

is the

active

electrodes If, at an nctizle electrode, oxygen is reduced to hydrogen peroxide which then disproportionates, it should be possible to trap hydrogen pero,xide before disproportionation is complete_ This possibility was realized by addition of Ti(IV) to the electrolysis cell before the recording of chronopotentiograms. Another experiment with the large platinum-gauze electrode was performed. The gauze electrode was activated by the method stated above and was transferred to the annular cell which contained an osygen-saturated solution of 0.02 _M Ti(IV) sulfate in I F sulfuric acid_ An oxygen-reduction chronopotentiog-ram was recorded. The electrolysis was internzpted at approximately +0_15 V z’s_ S.C.E. because, as shown below (curve 2, Fig. 6). the Ti(IV)-peroxide complex which is formed gives a

I

I

I

TIME

Fig. 6. Chronopotentiometric evidence for hvdrogen peroxide as unstable intermediate in osygen reduction at an n&-ive platinum electrode in 2 F HzSOA. Each curve was recorded at a current density of 737 PA/cm”. (I), Oxygen reduction in the presence of 0.1 AI Ti(IV) : (z), reduction of the Ti(IV) -peroxide complex in oxygen-free solution_ (3). normal oxygen-reduction chronopotentiogram; (4). reduction of Ti(IV) in oxygen-free soIution_

cathodic wave near -0.1 V ‘us. S.C.E. On the basis of the spectrophotometric method of analysis mentioned above, the yield of hydrogen peroxide was 5_7%_ Obviously, this yield is significantly larger than those obtained with active electrodes when .Ti(IV) was added to the solution a$er the electrolysis. W&&n the concentration of Ti(IV) was increased to. 0.x M, the hydrogen peroxide yield was S.So/o. Since 0.1 M is close to the solubility limit of Ti(IV) sulfate in I F sulfuic acid, a study of the trend of higher hydrogen peroxide yields with increasing Ti(IV) conc&tration was not extended. Further evidence .of the formation of hydrogen peroxide as an -unstable intermediate in .oTy&n rtiductioti at act&e electrodes was obtained from a series of conir&tio&l chronopotkntitimetrid experiments. The r&ults of these experiments J; EZiclroanaZ.

C?zem..

10 (1965)

306-318

316

D. G. PETERS,

R. A. MITCHELL

are shown in Fig. 6. Each chronopotentiogram was recorded after the following (anodization for zo set at a procedure : a platinum-wire electrode X\YLSacti\-ared current density of 737 ,u_A,!cm’ folio\\-ed b; cathndization at the same current density and the solution xvas stirred for approsljust to the hydrogen c~~~lution potential) matel>- 30 set and aXowed to become quiescent for at least one minute. Curve I of Fig_ 6 was obtained for os>-gen-saturated 2 F sulfuric acid containing 0.1 &1 Ti(IV). The first wave of curl-e I at +o.qs V z’s_ S_C_E. corresponds to the reduction of oxygen at the a&\-e electrode_ The second U-ELQ of curve I at -0-1 V z-s. S_C.E_ is due to the reduction of the Ti(IV)-peroxide comples r‘ormed by reaction of Ti(IV) with some of the hydrogen peroxide produced as an unstable intermediate in the reduction of oxygen. Evidence that this interpretation is correct is provided by the remaining three chronopotentiograms in Fig. 6_ Curve 3 is the cathodic chronopotentiograrn recorded for an oxygen-free solution of O-I 116 Ti(IV), o-005 F hydrogen peroxide, and z F sulfuric acid; essentially all the hydrogen peroxide was complexed by l?(W). Curve 4. obtained under the szune conditions as curve 2 except that h3-drogen peroxide was nbselzt, confirms that Ti(IV) by itself is not reduced at a potential more positive than -0.2 V ZIS.S.C.E. From curves 2 and 4, we conclude that the wave at -0.1 V vs. S.C.E_ is due to the reduction of the Ti(IV)-peroxide complex because it is a necessary and sufficient requirement that both species be present. Therefore, these experiments suggest that the reduction of ox>-gen at an active electrode in the presence of Ti(IV) yields hvdrogen peroxide, most of which disproportionates but some of which, being complexed by Ti(IV), causes the wave at -0.1 V ‘JS. S.C.E. It should be noted that the transition time for the oxygen wave, which occurs at f-0.45 is present (curve I) than when Ti(IV) is absent V vs. S.C.E., is shorter when Ti(IV) (curve 3)_ We attribute this to the fact that less oxygen is regenerated by dispropo-rkionation of hydrogen peroxide when Ti(IV) is present. The Ti(IV) -peroxide wave (curve z, Fig. 6) has several peculiarities, mentioned only briefly here, which will be investigated further. Beside the fact that the reaction responsible for the wave is unknown, the wave itself is not caused by a normal diffusion-controlled process. In fact, under conditions corresponding to those for which curve 2 was recorded, the Ti(IV) -peroxide wave decreases in size and vanishes after five or six successive cathodic trials; it reappears following re-activation (anodization and cathodization) of the electrode, whence the same sequence is repeated_ Moreover, activation of the electrode in the presence of the Ti(IV)-peroxide complex is not necessary for the appearance of the Ti(IV)-peroxide wave; identical results are obtained if the electrode is activated in pure sulfuric acid immediately before use in the Ti(IV) -peroxide solution. Another interesting point is that reduction of a platinum oxide film in air-free Ti(IV) solution produces the characteristic Ti(IV) peroxide wave. The latter observation suggests that the platinum oxide film consists of at least some adsorbed ox_)gen which is reduced to hydrogen peroxide or, perhaps, of a peroxide compound of platinum_ Aside from the evidence presented here, several controlled-potential coulometric studies have demonstrated the formation of hydrogen peroxide from reduction of oxygen at active platinum electrodes_ SAWYER AND INTERRANTE' measured hydrogen peroxide yields between 5 and 10% in-the controlled-potential reduction-of oxygen at an oxidizes electrode in 0.1 F potassium sulfate medium (pH 2). Under conditions employed by these workers, an ox3ized electrode is reduced almost im-

REDUCTIOS

OF 02 ,\T Pt-ELECTRODES

317

IS H~SOA-MEDI-+

mediately

to an nct~ve electrode”‘. Results (1o-2o~/~ yields) very SAWYER _am IXTERRQXE were obtained with an nctz& sulfuric acid”. EZOZA\X~AG found ho-I syb yields of hydrogen peroxide less controlled-potential reduction of osygen at an ohliz& electrode reported

by

similar to those electrode in I F for the more-orin 0.5 I; sulfuric

acid. Xlthouglr the electrode surface conditions in these experiments were less welldefined than in chronopotentiometry, all of these yields are small probably- because of the surface-catalyzed disproportionation of hydrogen peroside at an active electrode_

A discussion

of the mechanism

for oxygen

reduction

should

be qualified

by the

statement that only the overall products and important intermediate species can be identified with certaint>-. Details of the exact sequence of electron-transfer steps and intervening chemical reactions remain a subject for speculationOn the basis of the results reported above. a relatr\vely simple picture of the gross aspects of oxygen reduction can be proposed We postulate that, regardless of lvhether an active or n,aed platinum electrode is employed, oxygen undergoes electrochemical reduction to hydrogen peroside in a two-electron, diffusion-controlled step. e=HsOz

01,+2H++2

Although perhaps only a happenstance, at an active electrode (curve 2, Fig. 2) the reduction of osygen in I F sulfuric acid occurs at a potential very close to the standard potential (+o.++ V 2s. S.C.E.) for the osygen-hydrogen peroxide couple_ However, as the electrode ages, the olrerpotential for reduction of oxygen to hydrogen peroxide increases by at least 400 mV. An ectize platinum electrode possesses the ability to catalyze the disproportionation of hydrogen pero_xide to water and oxygen. HzOr? +

H-20+3

OS

M7e believe that this surface-catalyzed decomposition of hydrogen pero,xide occurs very rapidly during the electrochemical reduction of oxygen to hydrogen pero_xide at an nctive electrode. Consequently, the combination of these two reactions results in the overall diffusion-controlled, four-electron reduction of oxygen to water. Any process which acts to diminish or prevent the disproportionation of hydrogen peroside will cause the apparent z-value for oxygen reduction to decrease below four. Our results indicate that, as an active electrode is allowed to stand in oxygen-saturated sulfuric acid medium, the rate of disproportionation of hydrogen peroxide becomes sIower and slower until the apparent wvalue for oxygen reduction becomes two. The phenomenon of agilzg, whereby an nctize electrode ostensibly loses its deposit of finely-divided metal as well as the ability to catalyze hydrogen peroxide disproportionation, is subject to a variety of explanations. As possible alternatives, we include the following: (i) adsorption of trace impurities (poisoning) causes de-activation of the finely-divided metal; (ii) a physical recrystallization process occurs, as suggested by FRENCH AND KUWANA l5; and (iii) oxygen chemically oxidizes the finely-divided metal to platinum oxide and the oxide dissolves in the sulfuric acid mediurnr. Evidence can be found to support each possibility; however, no clear picture of the mechanism of aging exists at presentWork is being continued on a number of problems related to o_xygen reduction at platinum

cathodes.

These

include

the platinum-oxygen

reaction,

the phenomenon

/. Electromzal.Chetn., IO (1965) 306-318

D.

31s of aging, about

and voltammetric

the mechanism

and coulometric

of oxygen

reduction

reduction appears

G.

PETERS,

of osygen.

R.

Further

A.

MITCHELL

speculation

premature.

Chronopotentiograms for the reduction of osygen at an irct~ve platinum cathode (one which has a fresh deposit of finely-divided platinum metal on its surface) in sulfuric acid media show a single w-ave corresponding to the diffusion-controlled, fourelectron reduction to water. Howe\-er, the electrode ages rapidly upon standing in the solution, the overpotential for oxygen reduction increasing and the >z-value for the reduction becoming two. Direct chemical analysis of the cell solution following a single chronopotentiometric experiment has confirmed that hydrogen peroxide is not a stable product of the reduction at activa electrodes, but that hydrogen peroxide is quantitatively formed in the reduction at aged electrodes. Additional experiments with Ti(IV) present in the cell solution as a scavenger for hydrogen peroxide demonstrated that hydrogen peroxide is an unstable intermediate in the reduction of oxygen at active electrodes. We have concluded that, at active electrodes, oxygen is reduced to hydrogen as rapidly as it is formed, to water and more peroxide which disproportionates, oxygen at the electrode surface. As the electrode ages, it loses the ability to catalyze this disproportionation of hydrogen peroside and the reduction proceeds just to hydrogen peroxide. At aged electrodes the reduction of hydrogen peroxide to water does not occur due to the increased over-potential for this reaction. REFERENCES I J. J_ LIXGAXE, D. T- SAXVTER 3 D. T. SAXTER

?

4 A.

5 A. 6 7 S g

IO II

12 13 14 x5 x6

17 IS Ig_

20

zr

22 1.

A.

KOZAWA. KOZAWA. KOZAWA.

J_

_T_Electvoanal. Chem_. 2 (x961) 296. AXD I.- V. INTERRANTS, / EiectroanaZ. Chem., 2 (1961) AND R. J_ DAY, EZeclrochim. Acta, 8 (1963) 589 EZec~roc7zem. SOL JapaJr. 31 (x963) 183.

SOL Japan. 31 (1963) 61s. Gem., S (1964) 20. C. &-sax AXD D. M. KIXG. AmaZ. Them., 34 (1962) 362. C. AXSOK AXD J- J- LINGAXE. J. dm. Chenz. SOL, 79 (1957) qgoI_ A. LAITUSTN -4x.-D C. G. ENKE, J- Elecluochem_ SOL, 107 (1960) 773_ W. F-BERG, C. G. ENKE AND C. E. BRICKER.]. EbctrochemSot-, S. MA~ELL -D S. H. LANCER,]. Electrochzm. SOL, III (1~64) 438. J_ L.TNGAWE,~_ EZeclroanaZ. Chem_, I (1960) 37g_ C. Axso~.J. Amz_Chem.Soc.. 81 (1959) 1554_

310.

J_ Eleclrochenx J_ Eledroanal.

F. F_ H. S_ IIO 1963) 826. J. J_ F_ F. C. ANSON.~~~~~T CRem_, 33 (Ig6I) 934. W. G. FRENCH P-ND T. KuwAxA, J_ Phys. Chem., 65 (1964) IzFge A. SETT)EU, SoZzzcbiZitiesof Ixopga&c alrd Metal Organic Compo~tnds. D. Van Nostrand Co., Inc.. New York, 1940. J_ J_ LINGANE AND B. A. LOVERIDGE. J- Am. Chem. SOL. 72 (1950) 438. D_- G_ PETERS XND J- J- LINGANE. J- EZec#roanaZ. Chem.. 2 (1961) I_ T_ R. BLACEEBURN AND J- J_ LIXGANE, J. Electroanal. Chem_, 5 (x963) 216. J_ J_LINGANE AND P. J. L.INGANE.J. EZecfroanaZ. Chem., 5 (1963) 411~ H. A_LAITINEN AND I-M_ KOLTIXOFF,]. Phys.C?tem_. 45 (1941) 1061. D. G. PETE~~.unpublishedresults. _ EZectroanaZ_

Chem-.

IO

(1965)

306-3~8