J. R. GOLDSTEIN AND A. C. 6. TSEUNG
3646 hanIg have sta,ted that polymerization of olefins occur during ferricyanide-.thiol reactions. Similarly TurneyZohas also Ahown that radicals are produced during oxidation of mercaptans. Acknowledgment, One of us
(R. I(.C.) is grateful to
the Council of Scientific and Industrial Research, New
Delhi, for the award of a fellowship. Thanks are due to Evans Chemetics Inc. for the gift sample of 2-dimethylaminoethanethiol hydrochloride. (19) I. M. Kolthoff and E. J. Meehan, J . Polym. Sci.,11,71 (1953). (20) T. A. Turney, “Oxidation Mechanism,” Butterworths, London, 1965.
xygen Reduction on Graphite I Cobalt-Iron Oxide Electro eterogeneous Chemical Decomposition of H,O, by $1-R. Goldstein and A. C. C. Tseung* Demaytment of Chemistry, The City University, London, E.C.1, England
(Received March $0,1973)
Publication costs assisted by The American Chemical Society
The oxygen reduction reaction can be treated as a catalytic regenerative process; in alkaline solution, oxygen is initially reduced to give hydrogen peroxide, O2 HzO 2e = I l 0 ~ - OH-, which is then catalytically decomposed, 2H02- = 0 2 20H-, yielding oxygen for recycling. I n the present analysis of this system, exprewions have been derived predicting a higher open circuit voltage and exchange current density for an oxygen electrode incorporating the regeneration step, over that with no regeneration; both increases are mathematically related to the rate constant for hydrogen peroxide decomposition in the regeneration step. These effects have been studied using graphitelcobalt-iron oxide electrodes as model systems for oxygen re& ction by catalytic regeneration. On graphite, the 02-H02 - reaction proceeds reversibly, while cobaltiron oxides are good hydrogen peroxide catalysts possessing negligible oxygen reduction activity. Porous hydrophobic graphiteloxide electrodes were characterized in terms of diffusion-free pseudoexchange current density parameters. Incorporation of cobalt-iron oxides of known hydrogen peroxide decomposition activity into graphite electrodes produced increases in both open circuit voltage and pseudoexchange current density (io*) in agreement with theoretical predictions. Furthermore, the theory also indicated that only limited performance improvements are possible with oxygen electrodes operating solely according to the regenerative route. The maximum possible enhancement for graphitelcobalt-iron oxide electrodes is a doubling of io*, and the improvement in performance from an oxide of given hydrogen peroxide decomposition activity falls off with increasing io*.
+
+
1. ~ ~ t ~ o d u ~ t ~ ~ ~ The electrochemical reduction of oxygen in alkaline solution should ideally proceed according t o the equation 0 2
-/- 2WzO
+ 4e = 40H-
(1)
at a standard reduction potential of 0.401 V.I On most material&, howver, reaction 1 is slow and irreversible at room temperature. This causes performance limitations in power sources such as fuel cells and metal-air batteries, in which oxygen reduction is the cathodic reaction.2 A considerable research effort has been directed at developing ,an oxygen electrode at which reaction 1 might occur more etliciently, thereby yielding acceptably high-current densities at potentials close to the The Journal of Physical Chemistrv, Val. 76, No. $4, 1979
+
+
theoretical. A suitable electrode must be structure optimized, so that the three-phase reaction 1 can proceed with minimal mass-transport restrictions; the electrode must also incorporate electrocatalysts to increase the rate of the electrochemical reaction. Hitherto, the most promising results have been obtained for electrocatalysts such as platinum black and silver, but even with these costly materials, the electrode potential a t practical current densities falls well below the reversible value for reaction 1. One approach t o the oxygen electrode problem is b a e d on the fact that, for many electrocatalysts, oxygen reduction occurs initially by a two-electron pro(1) w . H. Latimer, “Oxidation Potentials,” 2nd ed, Prentice-Hall, Englewood Cliffs, N. J., 1952. (2) W. Vielstich, “Fuel Cells,” Wiley, 1970, pp 119, 327, 468.
OzREDUCTION ON GRAPHITE/ COBALT-IRON OXIDEELECTRODES cess.1 This produces hydrogen peroxide (H20z), present in alkaline solution as the perhydroxyl ion, HOZ-. 0 2
i- HZ0
4-2e
=
HOz-
+ OH-
(2)
React]on 2 occurs rapidly, but its standard reduction potential, -0.076 V,’ lies very much below the corresponding potendial for reaction 1. However, the HOzniay be chemically decomposed at the electrocatalyst surface 2H02-
=
0 2
+ 20H-
(3)
If the product Dxygen is completely recycled via reaction 2, then the overall stoichiometry of ( 2 ) and (3) approaches that 3f the four-electron process (1). Xoreover, the Nernst equation for reaction 2 at 25”, requires
where E is the electrode potential vs. a standard hydrogen electrode (SHE), and the square brackets represent activities. Assuming unit activities for 0 2 , €120, and OH-, then if [K02--]< 1, E’ rises above -0.076 V us. SHE. A low value for [HOz-] would result if the electrocatalyst were markedly active in HzOz decomposition. In particular, if HO2- could be maintained at conceiitrations as low as 10-l6 M , E would then be equal to 0.4 V, the standard reduction potential for reaction I .4 Hence, by setting up an oxygen electrode system based primarily on fast reactions 2 and 3, it might in principle be pxsible to approximate to the overall stoichiometry and high potential of (l), without its int;rinsic ~luggishness.~An important restriction would be the exclusion of any further chemical or electrochemical reactions. Thus, on platinum black and silver electrocatalysts, both reactions 2 and 3 proceed rapidly, yet reversible potentials are not attained. On such materials, however, many other reactions can simultaneously occur during oxygen reduction. These may either be largely irreversible, as in the case of furiher electrochemical reduction of HOz-
+
€302- HZO
+ 2e = 3OH-
(5)
or involve species with low standard reduction potentials,‘ for instance, the Pt(OH)z/Pt couple (0.15 V).l Tn both cases the working electrode potential mill be below 0.401 V vs. SHE. From the viewpoint of electrode kinetics, an oxygen ielectrode operating solely according to reactions 2 and 3 could be described in terms of an electrochemical reaction with a coupled heterogeneous regenerative (catalytic) reaction.% In the literature, however, there is little quantitative work available on such an oxygen electrode system. Conventional oxygen electrodes are empirically optimized, with a complex and ill-defined elcctroca talyst composition on which many reac-
3647
tions can occur simultaneously. Although the importance of an electrocatalyst also being a good H202decomposition catalyst is frequently stressed13n10there has been no quantitative approach to correlate the catalyst activity toward Hz02with the overall electrode performance. I n the present paper, the properties of a composite electrocatalyst system, operating essentially via reactions 2 and 3 alone, are described. A mathematical correlation is developed relating electrochemical performance to HeOzdecomposition activity, and is verified experimentally. The correlation is finally used to evaluate the Hz02 decomposition route as a means of achieving reversible oxygen electrode behavior.
2. Theory McIntyre,*following earlier work by Delahay and coworkers1Jr12 and Koutecky,13has considered the kinetics of an electrochemical reaction such as (a),coupled with a regenerative heterogeneous catalytic process such as (3)) occurring on or close to an electrode surface. An extended form of his analysis, applicable to the case in which reaction 3 is first order with respect t o HOz-, is presented here. In schematic form the overall electrode process may be represented as follows. AE
rB,
+ ne
-%A,
kE
A
BE
(6)
k-E
+ other products
(7)
Thus, equation 6 represents the electrochemical slow step, in which the oxidant A is electroreduced to the reductant B with the passage of 71 electrons. Subscripts E refer to an electrochemical process, with ICE and LEthe formal heterogeneous electrochemical rate constants for the forward and backward reactions, respectively. Equation 7 represents the coupled heterogeneous chemical regeneration reaction, in which the product B from the electrochemical reaction is decomposed, regenerating A, which is recycled in eq 6. In (3) W. G. Berl, Trans. Electrochem. SOC.,83, 253 (1943), (4) A. B. Hart and G. J. Womack, “Fuel Cells: Theory and Application,” Chapman and Hall, London, 1967, p 73. (5) J. R. Goldstein and A. C. C. Tseung, A’ature (London), 222, 869 (1969). (6) K. J. Vetter, “Electrochemical Kinetics,” Academic Press, New York, N. Y., 1967, p 639. (7) T. P. Hoar, Proc. Roy. Soc., Ser. A , 142, 628 (1933). (8) J. D. E. McIntyre, J. l‘hys. Chem., 71, 1196 (1967). (9) H. M . Cota, J. Katan, 14. Chin, and F. J. Schoenweis, Xiature (London), 203, 1281 (1964). (10) K . V. Kordesch in “Handbook of Fuel Cell Technology,” C. Berger, Ed., Prentice-Hall, Englewood Cliffs, N . J., 1968, p 361. (11) P. Delahay and G. L. Stiehl, J . Amer. Chpm. Soc., 74, 3500 (1952). (12) P. Delahay, C. C. ;Mattax, and T . B, Rerzins, abid., 76, 5319 (1954). (13) J. Koutecky, Collect. Czech. Chem. Commun 18, 311, 611 (1 953).
The Journal of Physical Chemistry, Vol. 7 6 , .Yo. 24, 1972
J. R. GOLDSTEIN AND A. C. C. TSEUNG
3648 eq 7 , subscripts c are used to refer to a chemical step, and the heterogeneous rate constant kc is defined according to the formation rate of A,
A
+ {r/(r - @]ne-+ (1
- $ ) { r / ( r - 6 ) f B + other products
(16)
Hence, for a diffusional flux SA, the cathodic current density is given by
i/nF where t is time, T is the stoichiometric ratio, moles of %,/moles of iL, of reaction 7, and s is the reaction (order with respect t o B,. The flux of A, J A ,is the net gain of A from both electrochemical and chemical sources
[A] and [B] are the solution concentrations of A and B, respectively, close to the electrode surface; the E and c subscripts in these terms need no longer be specified. Similarly, J B , the flux of B, is the net gain of B.
J B -- -d[B1 I-. dt
ks[A]
- ~ - E [ B ]- rkc[BIs
(IO)
Equations 9 and 10 give, on summing the fluxes of A and B, the equation
JA
-+- ,In +
(T
- 1)7~,[B]~ 0
(11)
The cathodic current density i is given by i/nP
=
IcE[A] - ~ - E [ B ]
(12)
Using eq 9 and 10 to eliminate A from eq 12 i/nF
=-
--JA
+ icc[BIs= J B + rk,[BIs
ffux of 13 from chemical decomposition flux of 13 produced electrochemically
(17)
From eq 9
i / n F = {r/(r - 6 ~ ) ) ( k E / A l- k--E[BI - I C ~ [ B ] ~ \ (18) or, neglecting the back electrochemical reaction
In McIntyre’s analysis, predictions were made as to the limiting current densities associated with infinitely fast regeneration rates. The present analysis, however, considers the contribution of the regeneration stcp t o the overall electrode current density i. An electrode system represented by eq 6 and 7 . should, in the absence of diffusion control, obey a Tafel relationship ?a
=
2.3RT7(10g i -- iogjoj anF
-.
with va, the activation polarization; R , the gas constant; T , the absolute temperature; a,the transfer coefficient; F , the Faraday; and jo,the exchange current density for the system. In the case of a negligibly slow regeneration reaction ( k , = 0 in eq 7), ,iofalls to a lower value io, the exchange current density for the slow electron transfer reaction 6 alone. Kow, substituting eq 20 into eq 15 with the condition va = 0, gives
But,, from the stoichiometry of 16
or, on rearranging
(14)
Whereupon
and (I - 0) is the fraction of the electrolysis product which escapes into the solution by mass transport, thereby avoiding thch regeneration process. By considering each mole of A supplied to the electrode surface by either rnass transport or regeneration, and summing all products over the course of the reaction as a geometric series, McIntyre showed the overall electrode reaction t o be The Journal of Physical Chemistry, Val. 76, N o . 24, 1,973
- $){JA
(13)
Equations 11 and 13 constitute the general boundary conditions for %hi3system. McIntyre defined $, the regeneration rate fraction, as that Fractiox of the electrolysis product BE which is consumed in the regenerative cycle after each occiirrence of the eleclrolysis reaction. 0 is accordingly a measure of the regenerative efficiency of the coupled catalytic process, and may be defined by the fraction 6
= --{r/(r
Expression 24 gives k,, the rate constant for the regeneration step, in terms of j o , the cxchange current density for the system, 8, the regeneration rate fraction, and [A], the oxidant concentration in the electrode vicinity. This expression may be greatly simplified, particularly in its application to experimental results, by employing an independent expression for 0. Thus, the increase in exchange current density from io, when there is no regeneration, t o j o , when the regeneration step is effective, must arise entirely from the product h recycled via eq 7. Inspection of eq 14 and 15, defining 6, should yield an expression for 8 in terms of i o
O2 REDUCTION ON GRAPHITE] COBALT-IRON OXIDEELECTRODES and io. The difference term ( j ~ - io) represents the flux of A arising from the regeneration step; since each mole of A arises from r moles of B, the numerator in eq 14 may be written r(j0 - io). Proportionally, the denominator is given by j o ,whereupon
(25) 30
The two expresslons, (24) and ( 2 5 ) , relate mathematically the three parameters io, j,,and k, governing the electrode kinetics of the regenerative system. Furthermore, the boundary conditions for the regenerative effect are satisfied. For example, with no or ineffective catalysis, zo = jo and 6 = 0 (eq 2 3 . Substitution of 6 = 0 in eq 24 gives k, = 0 as required. The case of ideal catalysis, ff = 1, may be demonstrated for a regeneration reaction which is first order with respect to B (s == I), and in which the stoichiometric ratio, in eq 8, is 2 . Substitution of 6 = 1 and T = 2 in eq 25 gives j" = 2i0,while, with the added condition s 1= 1, eq 24 gives k , = a . Both these results arp conditione for ideal catalysis.* 14
3. Oxygen Electrode System I n order to subject the theory of Section 2 to experimentalLanalysis, it was necessary to set up a model oxygen electrode system functioning essentially via reactions 2 and 3 above A two-component electrocatalyst system, with reaction 2 alone occurring on one component and reaction 3 alone occurring on the other, seemed most convenient for study, since kinetic effects corresponding t o either process could easily be separated, and the relative contribution of the regeneration step to the overall process directly controlled. Fundamental oxygen electrode investigations are normally carried out with the test material in solid form, since this permits a simple kinetic treatment. The dual electrocatalyst system, however, implied an intimate mixtune of components, thereby closely resembling a supported catalyst system. I n addition a basic prerequisite for the model system, that reactions 2 and 3 be rapid (Section I), could only be achieved using high surface area materials. The solid electrode approach, involving low surface area test mecirnens of homogeneous composition, was thus precluded from these studies, and a method of evaluating high eurface eiectrocatalysis as fine powder mixtures was adopted. Eleatrocatalyst powders may be fabricated as porous electrodes and their performance assessed from currentvoltage curves under controlled conditions. These curves>may be expressed as 7-i plots, where 7 is the polari:sation from the observed electrode open circuit voltage (ocv), and i is the electrode current density (current per unit geometric surface area). Generally, a is sLLbdividedinto activation ( q & ) ,mass transfer (vm), and ohmic (ax) components. While vm and qR are
3649
largely structure and electrolyte dependent, the qr component is a measure of the intrinsic activity of the electrocatalyst for accelerating the electrode reaction. Consequently, to assess this intrinsic activity, mass transfer and ohmic effects must be corrected for in an electrode polarization curve. This ideally leaves a purely activation-controlled qa-i plot, from which fundamental electrochemical parameters (such as the exchange current density for the reaction) may be obtained. With solid planar electrodes, it is relatively easy to establish pure activation dependence; a notable example is the use of such electrodes in a rotating disk configuration, where thc mass transfer effects are accurakely known. With porous systems, however, the electrode processes are regulated by a number of variables, such as the electrochemically active surface area, the electrode pore structure, and the aggregate size of the electrocatalyst. The equations governing the diffusion of electroactive species in such systems are most complex, and the elimination, particularly of a,,,, from the 7-i plots is at best empirical. However, Tantram and Tseung,14 in a study of oxygen reduction on high surface area graphite in alkaline solution, showed that diffusion effects in porous electrodes could be greatly minimized by close control of the electrode structure. In their investigation, porous hydrophobic graphite electrodes were s h o n~ Lo exhibit reproducible perforniance characteristics, and mass transfer effects only became significant at very high current densities. The graphite employed was reversible with respect to the HOz- formatiori r e a ~ t i o n yet ,~ was rather inert as an HZOZ decomposition ~rctalyst.'~ Neither was any participation from other electrochemical processes, such as reaction 5 , envisaged. By incorporating HzOB catalysts in such graphite-based porous electrodes, it might be possible to obtain relatively diffusion-free oxygen reduction characteristics. The selection of a suitable HZOZdecomposition catalyst posed a considerable problem, since the most active catalysts for reaction 3 , silver, platinum black, and palladium,15 also catalyze other electrochemical processes such as reaction 2 , thereby excluding them from this study (Section 1). A catalyst for reaction 3 of similar high activity, but otherwise electrochemically inert, was required. However, Cota, Katan, Chin, and Schoenweiss had reported that cobalt-iron spinel oxides were among the most active HzO2 decomposition catalysts in alkaline solution. These oxides bear the gcncral formula Co,Fe3--z04,where x, the composition variable, can take values between 0 and 3. By using the hydroxide coprecipitation technique of Saito, Sugihara, and Sato,16 cobalt-iron oxides could (14) A. D. 8. Tantram and A . C. C. Tseung, Sature (London), 221, 167 (1969). (15) W. C. Schumb, C. N. Satterfield, and R . N . Wentworth, "Hydrogen Feroxide," Reinhold, New York, N. Y., 1955. The Journal of Physical Chemistry, Vol. 76, No. 24,1972
J. R. GOLDSTEIN AND A. C. 6. T,QEUNG
3650 be prepared in high surface area form over the entire composition range of the system. The resulting powders, particularly within the composition range 1 < x < 3, exhibited activities for reaction 3 comparable to platinum black, when tested under controlled conditions. lq The cobalt--iron oxides exhibited properties well suited to the present study. The activity toward M20z was high and readily adjusted by variation in x. Thore was no apparent catalyst degradation even on prolonged exposure t o HLh, and the oxides were highly insolubie in alkaline solution by virtue of their preparation technique. Corrosion resistance at oxygen reduction potentials might wcll prove favorable. The oxides, in addition, possessed poor electrical conductivity, several orders of magnitude below that of the graphite, as shown by powder compression measurements under controlled conditmls. The likelihood of the oxides supporting any elec trochemical processes comparable to reaction 2 on the graphite was thus small. Finally, the microstructural teharacteristics of the oxide powders were suitably differcnt from those of the graphite, the graphite ponder' being considerably finer than the oxides, The electrode structure from graphite-oxide mechanical mixtures would hence be overwhelmingly the fine graphite, at least in the presence ~
x Section
~
e
~
4.1. Prepcration and Pretreatment of the Oxygen Electrode C'owponenls. A series of cobalt-iron oxides was prepared by the hydroxide coprecipitation technique, with x values of 0, 0.6, 1.0, 1.5, 2.0, 2.4, and 3.0 so as to cover the complete composition range of the spinel system. The preparation and characterization of these oxides ar mcribed e l s e ~ h e r e . ' ~The graphite as P. Acheson graphite, Type L.4. employedI4 1% This was pretrentd (800", 3 hr, 02-free Nz) so as to remove surfwe impurities, and stored in a COz-free atmosphere. 4.2 Activity o i Mntertals Toward HzOz. The activity of the graphile and oxides for H202 decomposition was evaluated gasometrically using the technique of Cota, et al q , 1 7 Since the kinetic results on decomposition aci;ivily were to be applied to the electrochemical situation of oxygen reduction, the catalysts were pretreated (300", 1 hrj 0s-free Nz)so as to approximate to the electrode fdxication conditions, and in addition, stabilizer-free P-I+& cas employed for the kinetic runs. Catalyst samples (50 me) were injected into a reaction mixture consisting of 5 N KOH (50 ml, made up from Analar grade pellets, €lopkin and Williams, Ltd.) and 5 ml of 4.0 (=t-O.Olj A4 Hz02 (Laporte stabilizerfree 35% H & z dilutod appropriately, and standardized The reaction was conducted at 25 f before 8.1" and atmospheric pressure, and the rate of oxygen evolution naoriitored over the complete reaction period. The Journal of f'h2sical Chemistry, V d . 76, N o . 24, 1972
The rate constant and reaction order were evaluated, and the dependence of the rate constant on the catalyst mass investigated. All runs were corrected for selfdecomposition of HzOz. 4.3 Electrode Fabrication. The electrode fabrication technique was a modification of that employed by Tantram and Tseung.I4 Weighed quantities of electrocatalyst powder(s) and aqueous TeAon dispersion (Fluon GPl, I.C.I., Ltd.) were slurried in methanol and ultrasonically dispersed for a 5.min period. The dispersion was transferred by painting onto a nickel gauze and after drying at 70" the electrodes were cured (300", 1 hr, G-free Nz). Electrocatalyst loadings were calculated from the gauze weight before and after fabrication, assuming that the components of the electrocatalyst-Teflon dispersion were transferred to the gauze in the same ratio as in the original mix. Electrodes were fabricated from graphite alone, for each oxide, and for graphite-oxide mixtures corresponding to each oxide composition, thereby covering the complete composition range of the spinel system. The elcctrodes were fabricated so as to exhibit the following compositions : graphite electrodes, 10 ( 10.3) mg C ~ I - graphite, ~ graphite to TeAon ratio = 3: 1; oxide electrodes, 10 (k0.3) mg cm -2 oxide, oxide to Teflon ratio = 3: 1; graphiteloxide electrodes, 10 i=k0.3) mg ~ m graphite, - ~ 10 (10.3) mg CIII-~ oxide, graphite to Teflon ratio = 3 : 1. ~ were chosen ~ The ~loadings and~ Teflon content to pro- ~ vide optimum hydrophobic properties and performance! reproducibility, and to minimize in-series electrode variations. Electrodes were also fabricated from graphite-inert filler oxide mixtures, in order to establish that the electrode performance was overwhelmingly graphite dependent. The oxide filler was ferric oxide prepared by hydroxide coprecipitation; ihis was as inert toward Hz02as the graphite, yet possessed similar microstructural characteristics to the oxides. 4.4 Electrode Testing. The electrodes, mounted a t the end of nickel foil, and of known geometrical surface area (ca. 1 cmz) were tested for oxygen reduction in 5 N KOH at 25" using the floating cell t e ~ h n i q u e . ' ~ In the cell, a gold foil counter electrode was used, and the reference electrode was a dynamic hydrogen electrode (DHE) ,la calibrated against a standard mercury \mercuric oxide electrode. The electrolyte was freshly prepared from KOH pellets and deionized, distilled water, and the working electrode compartment was swept with oxygen gas at about 5. a h , supplied from a cylinder without further purification. Electrode ocvs vs. DHE under oxygen were measured using a high-impedance (20 Rlohm) digital voltmeter (16) T. Sato, M. Sugihara, and M. Saito, Rer. Eke. Commun. Lab., 11, 26 (1963). (17) J. R. Goldstein and A. C. C. Tseung, submitted for publication. (18) J. Giner, J. Electrochem. Sac., 111, 376 (1964).
~
O2 REDUCTION ON GRAPHITE\ COBALT-IRON OXIDEELECTRODES
3651
-
Table I : Observed and Calculated Open Circuit Potentials for Graphite and GraphitelCobalb-Iron Oxide Oxygen Reduction Electrodesa
Electi*odt system
Graphite alone Graphitelcobaltiron oxides
5
Value of z in C0,Fea-~04 for added oxide
First-order rate constant for Ha02 decomposition per unit mass catalyst k f , m at 25' (f5%), sec-1 g.-1
Observed ocv us. SHE (25"), mV (f5)
6.7 x 0.0087 0 0308 0.117 0.0927 0.175 0.243 0.0899
100 160 165 175 190 190 200 170
0
0.6 1.0 1.5 2.0 2.4 3.0
Calculated steady-state HOz- concn,
Calculated ocv us. SHE (25O),
mol
mV (110)
3.7
x
80
10-7
3.1 x IO-* 8.1 x 2.1 x 10-10
I
140 160
180 175 180 185 175
2.7 x 10-lo 1.4 X 10-lo 1 . 0 x 10-1"
2.8
x
For data on electrode fabrication and test conditions see Section 4.
---
_ _ I -
Table PI : Tabulation of
Electrode system
Graphite alone Graphitel cobalt-. iron oxides
-
&,*, jo*, e, kf,m, and Tafel Slope Parameters (25")
Value of x Tafel slope in C0~Fes-~04 (ri;5'%0), mV for added per current oxide density decade
0 0.6 1.0 1.5 2.0 2.4 3.0
Mean pseudoexchange ourrent density
ltegeneration rate fraction
j a * (iSC/o),
220 220 225 240 225 240 240 235
mA om-2
1.08 (io*) as io*
1.42 1.70 1.60 1.87 1.84 1.72
5. Results 5.1. Activity of Calalysts Toward HzOz. The HzOz decomposition neaction on Acheson graphite, and on cobalt--iron clxides prepared by hydroxide coprecipitation, was found to be first order with respect to HzOz in alkaline solution. l7 First-order kinetics were also reported by Parravano for H20z decomposition by sint>ered cobalt -ison oxidesjZ0and were explained in terms of a rate-determining surface reaction (topocliemicall control) involving reactant-substrate orientations. The firat-order kinetics plot for H202 decomposition by ColkrOk a t 25" is shown in Figure 1, while the rate constant per unit mass catalyst kr,, at 25" for each of the materials tested is included in Tables I and
*
mA em-2
0
0 0
0
0,479 0,729 0.650 0.845 0.826 0.744
0.995 2.91 1.25 5.89 5.13 1.83
sec-1 g-1
Negligible 0 0067 (4.0308 0.111 0.0927 0.175 0,243 0.0899 __I
I _ -
(DVM, Solartron LM 1620). Steady-state polarization curves for oxygen reduction were then obtained using a potentiostat (Wenking 68TS10), and the curves were corrected for ohmic drop using the interrupter technique. l9 Corrosion scans were also carried out, with the working electrode and electrolyte equilibrated with On-freeITz.
jo
Electroohemieai parameter e 2-e
First-order rate constant for HaOz de~omposition per unit mass catalyst
11.
k f , m was found to be directly proportional to the catalyst mass, at least in the range 5-100 mg, and, as expected for first-order kinetics, was independent of ~ were the intial HzOe concentration. The J G F ,values subject to experimental error amounting to 5+5%. 5.6. Electrode Testing. 53.1. Oxygen Reduction on Graphite and Graphite /Cobalt--Iron Oxide Electrodes. After preliminary cathodic-anodic cycling to remove surface-adsorbed impurities and to achieve optimum wetting characteristics for the hydrophobic structures, reproducible performance levels were achieved for oxygen reduction 011 graphite and graphitelcobalt-iron oxide electrodes. Electrodes were least susceptible to irreversible effects, such as flooding by the electrolyte or hysteresis on cathodic-anodic cycling, when subjected to low-current, density drains ( 50 mV and 5 < i < 100 mA cmiZ~ A rtepresentative plot for the graphite/CoFezO* The Journal of Physical Chem&try, Vol. 76,N o . $4, 1972
Log i, mA cm-'
Pigure 4. Tafel plot for graphite\CoFeLh electrode for oxygen reduction in 5 N KOH at 25".
system is shown in Figure 4. The electrodes thus obeyed a Tafel relationship (Section 21, but the slopc and exchange current density (pseudoexchange current density) derived from q-log i plots x-ould be characteristic of porous electrodes. In Table II the Tafel slopes and pseudoexchange current densities for the various electrode systems are prcsented, the experimental error for each parameter being As%. Each value in the table is the mean of several determinations with successively fabricated electrodes. 5.6.2. Oxygen Reduction on Cobalt-Iron Oxide Electrodes. Cobalt-iron oxides tested alone exhibited considerable inertness for oxygen reduction. Thus a t -400 mTT lis. SHE, the cathode current densities observed for the various compositions were in the range 0.1-0.5 mA cm-%, nearly three orders of magnitude below the corresponding values in the presence of graphite. The oxide ocvs under oxygen were unstable and irreproducible, 150 ( * X i ) mV lis. SHE, in sharp contrast to the behavior of the oxides in the presence of graphite, when ocv fluctuations decreased to mV. 5.$.3. Corrosion Scans. The cobalt-iron oxide electrodes, when scanned under &free N2, gave negligible corrosion currents, and similar ocvs t o those under oxygen. Graphite electrodes, after preliminary cathodicanodic cycling to remove adsorbed oxygen, performed similarly.
0, REDUCTION
ON
C:I?APHITEI COBALT-IRON
OXIDEELECTRODES
The experimental work had established the suitability of the graphite\cobalt-iron oxide system for the present study in terms of (i) relative inertness of graphite for reaction 3 ; (ii) high activity of cobaltiron oxides for reaction 3; (iii) corrosion resistance and electrochemical inertness of cobalt-iron oxides alone during oxygen redualion; and (Kv) fabrication of electrodes with conb-olled quantities of graphite and oxides, showing reproducible performance for oxygen reduction. It was now possible to show that reaction 2 proceeds reversibly on the graphite, and to assess quantitatively the response of the graphite-oxide system to the regeneration step, in terms of the theory of Section 2 . 6.1. Observec' U F U S . In Table 1, the ocvs vs. SHE developed by the various systems are listed, together with the first-order rrtte constant per unit mass cat~ 1,3 2 0 2 decomposition at 25". There is a alyst, k ~ ,for broad correlation bet>weenthe ocv value for a given electrode and the activity for HzOz decomposition of the oxide it co:itains. Thus the higher the value of kt,, the higher is the ocv 8s. SHE. This suggested that the OCVE nere governed according to eq 2 by the steady-state corxentrstion of HQz- maintained by the catalyst at the electrode surface. Accordingl:y, the 3 - 1 0 2 - steady-state concentration maintained by each catalyst was calculated, and used to predict the o w from the Nernst eq 4. Consider the case of a typical graphite\CoFezOr electrode. The DVM has an impedance of 20 Mohm in the 2-V range used for meaE;urements us. UHF,. During an ocv meaout 1 V, the electrode current density is - 8 A cm-,, corresponding to an HQzproduction rate (eq 2) of 2.3 X mol H02- sec-l cm-2 for the ho-elee2,ron process. At the steady state the rate of formation of HQz- from this source may be equated with the rate of destruction of HOz- from the CoFez04 contained in the electrode (the decomposition from the graphite is negligible compared with this). For 10 mg of CoFez04In the electrode, and with 7cf ,m = 0.117sec-l g;-', the steady-state NOz- concentration is 2,85 X !id'. I n the Nernst eq 4, [O,] = 1, and [OH-] = 8 for 5 N KQN at 25°.21 Moreover ~
partial vapor pressure 5 N KOH a t 25" ___ ps,rtiaI vapor pressure water at 25"
which is 0.58.'" Vjsing [EOz--] = 2.15 X 10-'O M gives E = 180 InV us, SHE in eq 4. This is in reasonable agreement with the observed value of 175 mV us. SHE. Errors in predicted ocvs are likely to be large, due to uncertainty in the Kernst equation parameters, particularly [ H - I 0 2 - - ] . Thus, a variation of =t: 10% in [NQz-] produecs a shift of +20 mV in the ocv. The NO2- value is isrgely dependant on kf,m, itself subject
3653
to experimental error of i s % . Moreover, there is some doubt concerning the application of the full kf, m values to the electrochemical situation, due to Teflon screening of catalyst particles and escape of H0,-. However, these factors should not affect the relative order of ocv values in similarly fabricated electrodes. Table I1 lists the predicted ocvs vs. SHE for the several electrode systems. There i s fair agreement with observed ocvs within the stated error limits, providing confirmation of the reversibility of reaction 2 on Acheson graphite, and justifying the Nernst equation reasoning of Section 1. The quantitative treatment implies that the observed ocv is a function of the DVM impedance, and the amount of HzOz decomposition catalyst present in the electrode. With a higher impedance value, slightly higher ocvs are to be expected; however, the contribution of impurity depolarizing processes might then become more significant. Similarly, the use of a greater mass of HzOz decomposition catalyst should yield lower [HOz-] values, giving higher ocvs. In addition, for a typical graphiteloxide electrode, it is possible t o calculate the necessary kf,mvalue of the oxide component, such that the reversible potential for reaction 1 (0.4 V vs. SHE) might be achieved by depression of [SO,-] in reaction 2. Considering such an electrode, fabricated and tested as for the electrodes in Table 11, a value for [HO2-] of 6 X M gives E = 0.4 in eq 4; this would require kf,m= 4 X lo6 see-' g-' for the oxide at 25". WOKever, the most active cobalt-iron oxide tested, Coz 4!?eo 8 0 4 , has only kf,, = 0.243 sec-' g-Ia Consequently, several orders of magnitude increase in eatalyst activities are required to achieve ocvs close to the reversible potential for reaction 1, if the oniy reaction path is via eq 2 and 3. 6.2 Electrode Performance Characteristics. As discussed in Section 3 , the elimination of ohmic and mass transfer effects from electrode polarization curve-3 1' s an important step toward obtaining fundamental electrochemical parameters for an electrode system. However, a reliable estimation of mass transfer effects is often impossible for porous electrodes. The present studies with hydrophobic graphite-based systems were rather exceptional, in that no evidence of dlffusionlimited plateaux mas observed for current densities up to 2 A cm+ (Figure 3). Since mass transfer effects may be neglected provided i < O.li~,where ir, is the limiting current density of an electrode system,I9 it could be assumed that the resistance free V-z curves were largely diffusion free, at least for current densities below 200 mA Throughout this range the electrode behavior was (21) R. A. Robinson and 13. A. Stokes, Trans. Faraday Sooc., 45, 612 (1949). (22) "International Critical Tables," Vol. 3, E. W . Washburn, Ed., McGraw-Hill, New York, N. Y., 1928, p 373.
T h e Journal of Physical Chemistry, VoZ. 76, No. 1& 1971
J. R. GOLDSTEIN AND A. C. C . TSEUNG
3654 most reproducible and ohmic corrections small and most reliable. Moreover, the corresponding r-log i plots had thcir region of best linearity within this range. Consequentlg, Tafd slope and pseudoexchange current density parameters (Table 11) derived from such plots would be representative of pure activation control in the porous systems. Austin23 has considered the equations governing Tafel behavior in porous electrodes. In the absence of ohmic effects, for a system in which concentration gradients outsidla .:he electrode may be neglected, a Tafel relationship may be developed for a simple pore, thin-film model. The model assumes that linear diffusion occurs along pores whose diameter is negligible compared with t h e r length, and that the electrochemical reaction is confined to a thin film. At low- current densities, the qa-i relationship is given by
va
==
4.6RT anF
(log i
- log j o * )
This differs from the usual Tafel form (eq 20), in that a pseudoexchange zurrent density term jo*, characteristic of the porous electrode system, replaces the usual exchange current density term jo, characteristic of solid electrodes. In addition, the Tafel slope (coefficient of the log i term) for porous systems has twice its normal value. The doubled Tafel slope is caused by depletion of' reactant in the electrode interior, by ohmic effects within the electrode pores, or by a combination of kioth The obseived Tafel slopes in Table I1 fall in the range 220-240 mV per current density decade. If the doubled 'I'afel slope relationship (26) holds for the Teflon-bonded electrodes, the true Tafel slope on Acheson graphite is 110-120 mV. This is in good agreement with the Tafel slope result (110 mV) found by Yeager, KroLwe, and RaoZ4for oxygen reduction on smooth graphite in alkaline solution. In eq 26 a Tafel slope of 248 mV rtlsults by setting a = 0.5 (its usual value) and 7% = 1. This implies that in all graphitebased electrodes there is one rate-determining step involving a single electron t r a n s f ~ r irrespective , ~ ~ ~ ~ ~ of WpOzdecomposition catalyst additions. Table IT also lists the Icf ,m values for HzOz decomposition by the electrode components, together with the pseudoexchange current densities for the various systems A kcy feature of the results is that the lowest pseudoexchange current densities are shown for electrodes containing graphite alone. As cobalt-iron oxideP with su:cessively greater activity for HzOz decomposition are incorporated into the graphite electrodes, the pseudoetchange current density steadily increases. It is now possible t o employ the theoretical analysis of Section 2 . Thc porous electrodes obey the modified Tafeji eq 26, so real exchange current dcnsities in Section 2 may be replaced by the corresponding pseudoexchange current densities. The pseudoexchange curThe Journal oj Phyeical Chemistry, Vol. 76, N o . g.4, 197.9
rent density on graphite alone is hence referred to as io*; in the presence of appreciable regeneration this increases to a higher value, j o * . For cobalt-iron oxides involved in HzOz decomposition, the reaction order with respect to HOz- (9) is unity (Section 4.2), while the stoichiometric ratio (r) for the regeneration step is 2 (eq 3). In addition, the rate constant for HzOz decomposition, Lo, can be replaced by measured firstorder k f , m values, and n, defined by eq 6, is unity for eq 2.23r25 Equation 24 may be rewritten as
where 8, the regeneration rate fraction, from eq 25 is given by
In Table I1 the io" and j o * values are used to calculate values for 6' corresponding to each electrode system. The 0 values are employed to evaluate the electrochemical parameter (29) for each system, and in Figure 5, the calculated values for expression (29) are plotted vs. I%*,*, thereby testing the validity of eq 27. Figure 5 is subject to an error margin dependent on the individual variation in the ICf,m, io*, and jo* parameters. Although kinetic measurements determined kf?, to within j=5Oj,, a maximum accuracy of 10% N
E
0
U /'
,--. D I
O
_I_j
0.1
0.2
0.3
0.4
Figure 5 . Plot of the electrochemical parameter ( 8 / ( l - 8 ) ) { ( 2 - 0 ) / 2 } j o * for oxygen reduction a t Acheson graphite] cobalt-iron oxide electrodes vs. kf,,, the first-order rate constant per unit mass catalyst for H:02 decomposition by the oxide, at 2 5 " . (23) L. G. Austin, ref 10, pp 63, 184. (24) E. Yeager, P. Krouse, and K. V. Rao, Electrochim. Acta, 9, 1057 (1964). (25) R. Cornelissen and L. Gierst, J . Eleetroanal. Chem., 3, 219 (1962).
O2 REDUCTION os GRAPHITE( COBALT-IRON OXIDEELECTRODES was adopted for the electrochemical situation. This would allow for Teflon screening of catalyst surfaces in the various electrode compositions, and compensate for departures from the required catalyst loading levels. A *5% error in pseudoexchange current densities arises mainly from extrapolation errors on 7 log i plots, from variations in propcrties of successively fabricated electrodes (particularly the ocv value), and from inaccuracy in meisudng the electrode geometric surface area. The errcr margin of .t5% is remarkably small for exchange current measurements, and partly reflects the high revcrsibility of reaction 2 on the high surface area graphite. Thus, observed pseudoexchange current densities are several orders of magnitude above real exchange currextt densities measured for oxygen reduction at solid electrodes. Figure 5 if atxordingly constructed allowing =k 10% variation in k r , I n and .t5% in io* and jo*. The error limits resulting are Large because 13 is proportional to a small difference between two large quantities, and expression 29 contains a f12 term. However, the experimental points fall within the stated error limits, providing strong, wridcnce for the validity of eq 27. The graphitelcobalt- iron oxide system thus obeys the predictions of Seetion 2 . Support is also given to the viewpoint that fke Tafel regions for thc graphite-based electrodes are effectively under pure activation control, and that the e ectrochemical parameters in Table I1 are uniformly affected by the Teflon loading and electrode structure. The slope in 3 g u r e 5, by analogy to eq 27, is related to [A], the oxygcm concentration at the electrode surface, and possenses the value 1.2 X g mol This quantity cannot be given an absolute physical significance, sir ce it is dependent on other unknown factors, such as the fraction of catalyst mass actually involved in H 2 0 2 decomposition or the electrochemically active specific surface area of the graphite. The slope is presumably a characteristic of the Acheson graphitebased electrodes. The present malysis allows predictions t o be made as to the feasiblity of achieving high efficiency oxygen electrode performance by the use of the regenerative path. The oxide additions t o Acheson graphite produced changes in zo* ranging from a negligible increase (composition x = 0, the least active Hz02 catalyst in the series) to a 70% increase (composition x = 2.4, the most active catalyst). Thus, the practical limit for the system (100% increase in io*, from Section 2 ) has almost been achiex-ed with present catalysts, and high efficiency ha6 not been achieved for oxygen reduction. ES a differen, carbon base mere employed for the electrodes, posseasing a much higher io* value than Acheson graFhi1 e, the intrinsic performance level in the absence of R&? deconipoFition catalysts would be greatly improved. I n principle it would then be possible to improve this new zo* still further by incorpo-
3655
rating cobalt-iron oxides. Consider a carbon base with io* = 10 mA showing an order of magnitude increase over io* for Acheson graphite. If it is further assumed that the electrochemically active specific surface area of the carbon is similar to that of Acheson graphite, such that the [A] term remains unchanged at its value of 1.2 X 10-E g mol emw2,then the addition of 10 mg of oxide of composition x = 2.4 (h,,= 0.243 sec-' g-') would now only produce an incream in io*of 20y0 (from eq 27 and 28 by successive approximation). This is considerably below the 70% increase obtained for the Acheson graphite system with ioh" = 1 mA yet this oxide possesses the highest activity for HzOz decomposition. The analysis implies that high-performance oxygen electrodes cannot be achieved using the regenerative route alone. Thus. with a graphite or carbon base possessing an acceptably high-pseudo-exchange current density io* for oxygen reduction, any added catalyst would require a rate constant for H202deeomposition appreciably greater than can be obtained nith presently available materials, in order to significantly increase io*. 7. Conclusions
The Acheson graphitelcobalt-iron oxide system has been shown to be a suitablc model for the investigation of oxygen reduction solely by the €IO,- formation reaction 2 , and the regeneration reaction 3. Teflon-bonded hydrophobic electrodes, fabricated from graphiteoxide mechanical mixtures, may be characterized reproducibly in terms of a pseudoexchange current density parameter obtained from Tafel plots. These plots are essentially activation controlled and possess a doubled Tafel slope from chat observed nith solid graphite electrodes. The oxide additions to the graphite electrodes produce an increase in both ocv and the pseudoexchange current density (io") of the graphite electrode alone; these effects may be correlated with kf,,, the rate constant per unit mass for X I 2 0 2 decomposition by the oxide. However, only limited improvements in oxygen electrode efficiency are possible Jvith presently available catalysts using the regenerative route alone. The maximum possible effect in graphiteloxide electrodes is a doubling of io*, and the effect for an oxide of given k f falls off with increasing io*, Conventional carbon or graphite-based oxygen electrodes, catalyzed with platinum or silver, and exhibiting much higher performance levels than the uncatalyzed systems, presumably rely on nonregenerative oxygen roduction routes for their high efficiency. Further work is required to test the applicability of the mathematical analysis to other systems. Of particular interest are carbon-based regenerative systems in which the carbon has an appreciably different i o * from Acheson graphite. It is possible to use a similar The Journal of Physical Chemistqil, VQE.76, N o , 24, 1972
A. J READ
3656 approach to investigate oxygen reduction in acid solution. A suitjable system for study is the carbon-iron phthalocyanine system, which closely resembles the graphite/ cObslt-iron Oxide system in its mode Of Oxygen reduction.
Acknowledgment. This work was supported by the Science Research Council on Grant No. B/SR 4334. (26) H. Jahnke and M , Sc+jn,erg, proc. J . I&. Etudes pilesCombust,, Brussels, 1060, 60 (1969).
Dissolution of Copper in Weakly Acidic Solutions by A. J. Read Chemistry Division, Department of Scientific and Industrial Research, Gracefield, N e w Zealand
(Received Pebruary 7 , 1978)
The dissolution of copper in weakly acidic perchlorate solutions was measured over a range of pI3 (4.0-5.5), PO, (0.05-1.0 bar), and cupric ion concentrations (1 X 10-4-16 X M) a t constant ionic strength (0.05 M ) ~The method used allowed several measurements to be made under pseudo-zero-order conditions on a single piece of copper in situ. The order dependence on (H+) and PO$was found to be 1.0 and 0.25, respectively. Increasing the temperature from 16.9 to 45.0' increased the rate by a factor of 6.2. Addition of C1-- and Br- ions enhanced the dissolution rate while added F-, NOS-, S042-, and Zn2+had a negligible effect. These results are discussed in terms of a solid-state mechanism involving the formation, transport, and destruction of cation vacancies and positive holes.
Introduction Numerous publications have appeared dealing with the corrosior. of copper under a wide variety of conditions. l - I G However, few authors studied dissolution under conditions allied to natural mater systems.6b16l7 The reaction mechanism was not fully elucidated under acidic and is even less well understood in solutions conhaining low concentrations of complexing ions. The present work seeks t o improve the understanding of the dissolution reaction in weakly acidic media (pH range 4.0-5.5). The oxjde films, formed initially on copper single crystals in water, were shown to consist of cuprous oxide@-*O even though thermodynamic considerations indicated cupric oxide was the stable phasesz1 At room temperature, Lhese films grew more rapidly in water than in air.lR It seemed likely, therefore, that cuprous oxide would be an intermediate in the dissolution of copper in the present study. It Is often difieult to control all the variables for a heterogeneous reaction, to a suitable level of precision. I n the present P ; O I ' ~ a technique was used which allowed the precise control of reaction variables (e.y., PO,,pH, (Cu2+),m d ionic strength) and minimized the problem of .aeprodiicing characteristic metal surfaces. Further, ns relai,jve rate measurements are sufficient for a mechanistic ttud.y, no great importance ivas attached to thcb variation of rate with surface preparation. The Journal of Physical Chemistry, Vol. 76, N o . $4, 1072
Experimental Section The samples of copper used in the present investigation were of Analar copper foil for which an emission spectrographic analysis is given. (1) B. C. Y. Lu and W.F. Graydon, Can. J. Chem., 32, 153 (1954). (2) 5. R . Weeks and G. R . Hill, J. Electrochem. Soc., 103,203 (1956). (3) R. P.Russel and A. White, I n d . Eng. Chem., 19, 116 (1927). (4) G. H. Damon and R. C. Cross, ibid., 28, 231 (1936). (5) I. Cornet, E . A. Barrington, and G. U . Behrsing, J . Electrochem. SOC.,108, 947 (1961). (6) A. I. Kinevski, J . A p p l . Chem. U S S R , 28, 1088 (1955). (7) W. D. Robertson, V. F. Nole, W. H . Davenport, and F. P. Talboom, Jr., J . Electrochem. Soc., 105, 569 (1958).
(8) 2. Zembura and TV. Glodzinska, Rocz. Chem., 42, 1525 (1968). (9) R . Glauner, 2 . Phys. Chem., Abt. A, 142, 67 (1929). (10) T. Hurlen, Acta Chem. Scand., 15, 1239 (1961). (11) 6. R.Hill, J . Electrochem. SOC.,100, 345 (1953). (12) J. Halpern, ibid., 100, 421 (1953). (13) 2. Bzabo, J . Amer. Chem. Soc., 7'1, 1511 (1949). (14) W. Kats, Werkst. Korros. (Weinheim), 1, 393 (1950). (15) T. Hurlen, Acta Chem. Scand., 15, 1246 (1961). (16) D . J. G. Ives and A . E. Rawson, J . Electrochem. Xoc., 109, 452 (1962). (17) 2 . Zembura and W. Glodzinska, Rocz. Chem., 40, 911 (1966). (18) ,J. Kruger, J . Electrochem. Soc., 106, 847 (1959). (19) J. Kruger, ibid., 108, 503 (1961). (20) A . Goswami, I n d i a n J . P u r e A p p l . Phya., 7, 211 (1969). (21) M. Pourbaix, "Atlas of Electrochemical Equilibria," Pergamon Press, London, 1966.
