1720
Anal. Chem. 1082, 5 4 , 1720-1724
Effects of Media and Electrode Materials on the Electrochemical Reduction of Dioxygen Donald T. Sawyer,* Glalco Chlerlcato, Jr., Charles T. Angells, Edward J. Nannl, Jr., and Tohru Tsuchlya Department of Chemistry, UnlversiYy of Californla, Riverside, California 9252 1
Cycllc voltammetry has been used to measure the extent of superoxlde Ion solvatlon (and consequent shift In the redox couple) In acetonltrlle, dlmethylformpotentlal for the 02/02amlde, pyridine, and dlmethyl sulfoxide. A parallel study of the reverslblllty of the 02/01couple as a function of electrode materlal has been made for platinum, gold, mercury, and glassy carbon In acetonltrlle. The results Indicate that surface Interactions between superoxide Ion and metal-electrode surfaces occur. Additional studies of the reductlon of oxygen In the presence of protic substrates and metal salts confirm that Lewis aclds promote the disproportkmatlon of superoxide Ion. I n the presence of electrophlles (e.g., CCI, and esters) the reduction current for dioxygen Is enhanced because of rapid nucleophlllc reactlons by 4-, which result In regeneration of 02.
Although the electrochemical reduction of dioxygen to superoxide ion in aprotic solvents has been known since 1965 (1-5), the effects of media and electrode material have not been addressed in any systematic study. An early study (6) established that activated metal electrodes in aqueous media catalyze the reduction of O2to water at a reduction potential characteristic of the particular metal. In contrast, at passivated metal electrodes the reduction of 02,which is independent of electrode material and pH, yields H202via a one-electron mechanism with facile disproportionation of the 0,- intermediate (7).Three recent reviews discuss the redox chemistry and electrochemistry of dioxygen (8-10). The general view has been that the electrochemistry of O2 in aprotic media also is independent of media and electrode materials, and that only protic contaminants promote the rapid disproportionation of the 02-product to H202 and 0 2 . For a test of this conclusion the present study has been undertaken to compare the effects of various media and electrode materials on the reduction potential, reversibility, and overall electron stoichiometry for the cyclic voltammetry of dioxygen. EXPERIMENTAL SECTION Instrumentation. The cyclic voltammetric measurements were made with either a Princeton Applied Research Model 173/ 175/179 potentiostat/universal programmer/electronic integrator or a universal potentiostat/amperostat that was constructed from operational amplifiers (11). The electrochemical cell (Brinkman) included (a) a platinum-foil auxiliary electrode that was isolated from the bulk solution by a glass tube with a medium frit, (b) a Ag/AgCl reference electrode in a glass tube closed with a cracked sofbglaas tip that contained aqueous Me4NCl (concentration was adjusted to obtain a potential of 0.000 V vs. SCE) (12),(c) a luggin capillary that contained the bulk electrolyte solution and the reference electrode (the capillary tip was placed within 1mm of the working electrode surface), (c) inlet/outlet tubes for gas dispersion in the solution or over its surface, and (d) the working electrode. The latter was either a Beckman Model 39273 platinum-inlay electrode (area, 0.23 cm2)or one of the inlay electrodes (Pt,Au, and glassy carbon; apparent surface areas, 0.74 cm2) of the Beckman rotating disk electrode assembly; the Hg electrode was a classical glass U-tube (6 mm 0.d.) with a Pt-wire 0003-2700/82/0354-1720$01.25/0
connection (estimated area, 0.3 cm2). A Beckman vitreous carbon-inlay electrode (area, 0.20 cm2)also was used for some of the measurements. The surface for each of the solid electrodes was polished with Buehler No. 3 (0.05 pm) polishing alumina prior to each scan. The negative voltammetric scans were initiated at the rest potential for the electrolyte/electrode system and were at a scan rate of 0.1 V s-l, unless otherwise stated. Reagents and Solvents. The aprotic solvenb were obtained from Burdick and Jackson ("Distilled in Glass" grade), and each had a "lot analysis" for water: acetonitrile (MeCN) (0.007%), dimethylformamide (DMF) (0.010%),dimethyl sulfoxide (Me2SO) (0.020%), and pyridine (Pyr) (0.006%). All of the solutions contained 0.1 M tetraethylammonium perchlorate (TEAP) (Reagent Grade, G. F. Smith Chemical Co.) as the supporting electrolyte. All other materials were reagent grade and were used without further purification. The solubility of O2at 1atm pressure and 25 "C in the four aprotic solvents (0.1 M TEAP) was evaluated by controlled potential coulometric reduction of a known volume of 02-saturatedsolution in an electrochemical cell that had zero head space. The approximate solubilities are as follows: MeCN, 8.1 mM, DMF, 4.8 mM; Me2S0, 2.1 mM; Pyr, 4.9 mM. RESULTS The electron-transfer reduction of dioxygen is affected by the solution matrix (solvent, protic substrates, reactive substrates, and metal ions) and by the nature of the working electrode. Through the use of cyclic voltammetry the specific effects for each of these variables have been determined. Media Effects. Solvent. Figure 1 illustrates cyclic voltammograms for the reduction of dioxygen (1 atm) a t a vitreous carbon electrode in four aprotic solvents acetonitrile (MeCN), dimethyl sulfoxide (Me2SO),dimethylformamide (DMF), and pyridine (Pyr). The peak potentials for the dioxygen reduction process and for the reoxidation of the product species (superoxide ion) in the four solvents as well as analogous data for three other electrode materials are summarized in Table I. Although the peak separation (AE,) varies with solvent and with electrode material, the median potentials (Ep,c Ep,,)/2 are essentially independent of electrode material and provide a reasonable measure for the formal reduction potential ( E O ' ) for the 02/02couple. The average values for Eo' are summarized in Table I together with the value for an aqueous solvent at pH 7 (13). The variation in the peak-separation values (AE,)for the cyclic voltammetric data of Table I may be interpreted in terms of heterogeneous electron-transfer kinetics, but the more reasonable explanation is uncompensated resistance (especially for Pyr and MeCN) and surface reactions (especially for the metal electrodes). Protic Substrates. When dioxygen is reduced in the presence of an equimolar concentration of a strong acid (HClO$ in DMF (Figure 2), a new irreversible peak occurs at -0.13 V vs. SCE in addition to the regular quasi-reversible couple at -0.86 V. (In the absence of O2strong acids (HC104) in DMF exhibit a reversible one-electron couple at -0.37 V). The peak height a t -0.13 V increases linearly as the concentration of HC104 is increased up to a mole ratio of 4:l relative to O2 (equivalent to a two-electron reduction). This increase in peak height is at the expense of the reversible couple at
+
0 1982 American Chemical Society
ANALYTICAL CHEMISTRY, VOL. 54,NO. 11, SEPTEMBER 1982
1721
Table I. Redox Potentids for the Reduction of Dioxygen (1 atm) by Cyclic Voltammetry with Four Different Electrodes in Four Aprotic Solvents, (0.1 M Tetraethylammonium Perchlorate) A. Electrochemical Peak Potentials; Scan Rate, 0.1 V s-l; V vs. SCE
-solvent Me,SO DMF PYr MeCN
C
E P , ~ Ep,a -0.86 -0.70 -0.95 -0.77 -1.00 -0.76 -0.9B -0.78
Aep
0.16 0.18
0.24 0.17
Pt E P , ~ Ep,a -0.87 -0.70 -0.96 -0.76 -1.03 -0.76 -1.18 -0.60
AEp 0.17 0.20
0.27 0.56
E,, -0.90 -0.98 -1.00
-1.12
Au Ep,a -0.68 -0.78 --0.75 -0.65
Hg
AEP
Ep,c
Ep,a
AEp
-1.12
-0.61
0.51
0.22 0.20
0.25 0.47
B. Formal Reduction Potentials for the O,/O,- Couple (Molar Concentrations for 0, and 0,-)
E"',V vs. SCE
E"', V vs. SCE
a
H,Oa'
- 0.41
Me,SO DMF
-0.78 -0.86
PYr MeCN
-0.88 - 0.87
Reference 13. I
I +IO
I 00
I
1 -10
I
-20
E , V vs SCE
Flgure 2. Cyclic voltammograms for 4.8 mM 0,(at 1 atm) in DMF (0.1 M TEAP) (solid lines) and in the presence of an equimolar concentration of HCIO, and in the presence of a 4:l mole ratio of phenol (PhOH) (dashed lines): platlnum electrode, area 0.23 cm2; scan rate, 0.1 V S-1.
Table 11. Half-Peak Potentials and Electron Stoichiometry for the Reduction of 4.8 mM 0, in DMF (0.1 M TEAP) in the Presence of a !%FoldExcess of Acidic Substrate at a Platinum Electrode (Scan Rate, 0.1 V s-') Figure 1. Cyclic voltammograms for 0, (at 1 atm) In acetonitrlle (MeCN) (8.1mM), dlmethyl sulfoxide (Me2SO)(2.1mM), dlmethylformam& (DMF) (4.8mM), and pyridine (Pyr) (4.9mM) at a vitreous carbon inlay electrode (area, 0.20crnI2). All of the solutions contained 0.1 M TEAP as supporting electrolyte, and the scan rate was 0.1 V s-'.
-0.86 V. The addition of excess phenol to B dioxygen-DMF solution causes the one-electron process to become a twoelectron reduction to yield H202as the major product (Figure 2); the reverse scan indicates that phenoxide is formed during the reduction of 02.Similar effects are observed when other moderately weak acids are present in excess, but the addition of a 5-fold excess (relative to O2 concentration) of H 2 0 or 1-butanol does not affect the electrochemistry of O2 in DMF (Table 11). Figure 3 illustrates the effect of HCIOl (70%) on the electrochemistry of 0 2 in cicetonitrile. The initial peak a t +0.11 V vs. SCE is due to proton reduction to hydrogen atoms, which react with an O2a t the surface. Hence, with an adequate flux of hydrogen atoms the O2is reduced such that its
substrate 3,5-di-tertbutylcatechol phenol a-tocopherol H,O 1-butanol
-0.85
2
-0.88
-0.88
2 2
-0.90 -0.90 -0.90
1 1 -1
concentration at the electrode surface is zero and the regular 02/02peak is eliminated. Organic Substrates. Previous studies have demonstrated that superoxide ion is an effective nucleophile in aprotic media (14-17). Hence, the presence of alkyl halides or carbonyl compounds with effective leaving groups (acyl halides, esters, and anhydrides) (18-22)in an oxygen solution results in an apparent increase in the electron stoichiometry for the cyclic voltammogram and a decrease in the reverse peak height
1722
ANALYTICAL CHEMISTRY, VOL. 54, NO. 11, SEPTEMBER 1982 I
I
I
I
I
,k:
I
8 : 8 02:HC104,'1
8 r n M HCI04
I
I
I
1%
+Zn(H,O)it//
,I
:;
I':
E, V v s SCE
Flgurr 3. Cyclic voltammograms in MeCN (0.1 M TEAP) of 8.1 mM 0,(at 1 atm) (solid Ilne) and in the presence of 8 mM HCIO, (short dashed line); and of 1.6 mM 0,(air at 1 atm) in the presence of 8 mM HCIO, (long dashed line): platinum electrode, area 0.23 cm'; scan rate, 0.1 v s-1.
Table 111. Examples of Organic Substrates That Enhance the Voltammetric Peak Height for the Reduction of 0, via Post-Electron-Transfer Reactions by 0,in DMF (22, 23)(The Source of 0,- Is the Reduction Process 0,t e- 0,-) -+
A. CH,Cl CH,CI + 0,- -+ CH,-0-0. t C1k , 80 M-l s-l overall stoichiometry 2CH3C1 + 20,- + CH,-0-0-CH, + 2C1- + 0, electron stoichiometry 2CH,C1 + 0, + 2e(in the limit as u + O)(l CH,-0-0-CH, + 2C1-; 2e-/0, B. CC1, rate-controlling step CCl, + oz- c1,c-oo~+ C1k, 1300 M-' s-l overall stoichiometry CCl, + 60,- CO,,- t 4C1- + 40, electron stoichiometry CC1, t 20, + 6e- CO,,- + (in the limit as u 0 ) 4C1-; 3e-10, rate-controlling step
-+
I
I
0.0
-0.5
I -1.0
I -1.5
E, V v s SCE
Figure 4. Cyclic voltammograms In Me,SO (0.1 M TEAP) of 2.1 mM O2 (at 1 atm) (solid line) and in the presence of 50 mM Zn(DMU),(CiO,), (short (DMU = dimethylurea) (dotted line), 50 mM Zn(H,O),(CIO,), dashed line), and 50 mM Zn(bpy),(CIO,), (bpy = 2,2'-bipyridine) (ion dashed line): platinum electrode, area 0.23 cm2; scan rate, 0.1 V s- .
B
Table IV. Voltammetric Reduction of 2.1 mM 0, in the Presence of a 10- to 30-Fold Excess (Relative to 0, Concentration) of Metal Cations in Me,SO (0.1 M TEAP)a at a Pt Electrode (0.02 V s-') metalb Li+ ZnZ+ CdZ+ Fez+ Mn2+ COZ+
EP,W
V vs. SCE
n (e-/O,)
-0.80
1 1 2 2 2
-0.79 -0.67 -0.67 -0.52
-0.61 -0.66
4 4
-+
-+
-+
-f
C. CH,C( 0)OPh CH,C(O)OPh + 0; -+ CH,C(O)OO. + PhO- k , 160 MV1 overall stoichiometry 2CH3C(0)OPh + 40,- -+ 2CH,C(O)O- + 2Ph0- + 3 0 , electron stoichiometry 2CH,C(O)OPh + 0, + 4e- -* (in the limit as u -+ 0 ) 2CH3C(0)O- + 2Ph0-; 4e-10, a u represents the voltammetric scan rate at a platinum electrode (0.23 cmz)for the reduction of 0, in the presence of excess organic substrate.
rate-controlling step
(these substrates are not electroactive within the voltage range for 0 2 reduction). The extent of this effect is dependent upon (a) the concentration of the substrate, (b) the stoichiometry for the substrate-02- reaction, (c) the rate constant for the latter, and (d) the voltammetric scan rate. Illustrative examples are summarized in Table 111. When an excess of a reactive substrate such as CC14is present, the peak height for O2reduction is more than doubled at a scan rate of 0.1 V s-l (the process is totally irreversible) (22); the cyclic is similar
a Tetraethylammonium perchlorate. hexahydrated perchlorate salts.
Added as the
in appearance to that for the 4:l PhOH:02 system (Figure 2). Metal Cation Substrates. The electrochemical reduction of O2in aprotic media is dramatically changed by the presence of electroinactive metal cations. Figure 4 illustrates the effect of a 5-fold excess of Zn(H20)6(C104)2,Zn(dimethylurea)6(C10J2, and Zn(bip~r),(ClO~)~ on the cyclic voltammetry for O2in DMF at a platinum electrode. Prior to each reductive scan the electrode has been repolished; a second scan yields a much reduced peak current. In the presence of an excess concentration of Zn(I1) cations as weil as the cations of Cd(II), Fe(II), Mn(II), and Co(II), the reduction of O2is a totally irreversible process, and the electrodes (Pt, Au, and C) are passivated after the initial negative scan. For slow scan rates in the presence of 10- to 30-fold excess concentrations (relative to the O2 concentration) of several divalent metal cations, the potentials for 0, reduction are shifted to more positive values and the peak currents are increased by a factor of 2 or 4. Table IV summarizes the results for such experiments for a platinum electrode in Me2S0 (0.1 M TEAP). Apparently, these cations act as Lewis acids relative to the 02product species from the electron-transfer step and thereby cause the reduction process to be shifted to more positive potentials in the order: Fez+ > Mn2+> Co2+
ANALYTICAL CHEMISTRY, VOL. 54, NO. 11, SEPTEMBER 1982
disproportionates and that it is not electroactive in DMF at potentials less negative than that for the 02/Orcouple (-0.86 V vs. SCE). A reasonable pathway for the facile disproportionation is the formation of a dimer HO2. HO2. [HOOOOH] k > lo7 M-l s-l (2) H202+ O2
I
,-Ai,
1
1723
-
+
-
with subsequent dissociation to dioxygen and peroxide. To achieve a full two-electron peak height for the process of eq 1 requires a ratio of at least four HC104 molecules per O2 molecule in DMF. This results because HC104 protonates DMF and the diffusion coefficient for the latter is much smaller than that for 02.Furthermore, the flux of HDMF+ to the electrode must be twice that for O2 to achieve the second cycle of eq 1 with the products of eq 2. In the case of O2 reduction in the presence of HC104-MeCN (Figure 3), the primary process is represented by (26) HC104(H20)
OC
E. V
-
DI8CUSSION Reference to Table I confirms that the reduction potential for the 02/02couple shifts to more negative values as the solvating properties of the media decrease. The heat of hydration (-AHq) for gaseous 02-is 418 kJ (24),which is consistent with the unique strong solvation of anions by water. Hence, if the Eo’0,,o,-values for the 02/Orcouple are affected primarily by the degree of solvation of 0; (that is, the solvation energy for O2 is assunied to be small and about the same for the different solvents), then the relative solvation energies for 0 2 - are H 2 0 >> Me2S0 > DMF > Pyr MeCN. The cyclic voltammograrn of Figure 2 indicates that a new process occurs when O2 is reduced in the presence of a strong acid. For example in DMF
-
-
+ C1O4(Hz0)-
H.(Pt)
Ep,c= +0.11 V vs. SCE (3)
+
0 20 vs SCE
Zn2+ Cd2+> Li+ ‘TEA+ (tetraethylammonium cation from the supporting electrolyte, TEA(C10,)). On the basis of the electron stoichiometries of Table IV, the Lewis acidsuperoxide adducts (M2+-.(02-),)disproportionate to metal peroxides and dioxygen. In the case of Mn2+and Co2+their peroxides are unstable and disproportionate to metal oxides and dioxygen to yield an overall stoichiometry of four electrons per 02. Electrode Material Effects. Figure 5 illustrates the cyclic voltammograms for the reduction of 0 2 in MeCN (0.1 M TEAP) at glassy carbon, platinum, gold, and mercury electrodes. The peak potentials and peak separations for the reduction of O2 and reoxidation of 02-with these electrodes in MeCN, Me2S0, DMF, and Pyr are summarized in Table I. Clearly, the larger peak separations (.A&,)‘ for the metal electrodes, especially in MeCN, indicate that the apparent irreversibility for the 02/Orcouple is due to the reaction of O,, or its disproportionation product (OZ2-),with the metal surface. This effect is largest in MeCN because it is the poorest solvating agent for 02-,which is equivalent to minimizing its deactivation.
O2 + HC104 + e-
Pt
which is followed by a chemical surface reaction H*(Pt) 0 2 HO2. Pt
Figure 5. Cyclic voltammograms in MeCN (0.1 M TEAP) of 8.1 mM O2 (at 1 atm) at vitreous carbon, platinum, gold, and mercury electrodes: electrode areas, 0.74 cm2(except for mercury, 0.3 cm2);scan rate, 0.1 v s-’.
- -
-
+ e-
H02. + C104E,,, = -0.13 V
VS.
SCE (1)
Previous studies (25) have shown that the HO,. species rapidly
+
+
(4)
The H02. species disproportionate via the reaction of eq 2. When excess quantities of weak acids such as PhOH are present in the aprotic solvent, the potential for the O2 reduction is not significantly affected (Table 11). For example, in DMF
O2 + e-
-
E , , = -0.89 V vs. SCE
02-
(5)
However, the increase in the peak height relative to that for a one-electron process and the absence of a reverse peak confirms that post-electron-transfer reactions occur
02-
+ PhOH 2 HO2*
+ PhOHz02 + 02
+
+
HOz.
-
(6) (7)
The overall reaction for excess PhOH is
O2 + 2PhOH
+ 2e-
H202
+ 2Ph0-
(8)
with the presence of PhO- as a product confirmed by the anodic peak at +0.3 V vs. SCE in Figure 2. When excess quantities of substrates that are subject to nucleophilic attack by 02-are present, the reduction peak for O2 is enhanced in height and the process becomes irreversible because of post-electron-transfer chemical reactions. Several examples are summarized in Table I11 to illustrate the extent of the effect. The presence of metal cations also promotes post-electron-transfer chemistry with the 02-from the O2 reduction (5). The results of Figure 4 and Table IV can be rationalized on the basis of a Lewis acid-base process that promotes disproportionation to metal peroxide and 02.For example, with excess zinc ion relative to O2 in aprotic media Zn2+
-
+ 02-
-
Zn(02)+
Zn(02)++ 02- [Zn(O,),l
-
Zn02 + O2
(9)
(10)
Preliminary studies (27) by stopped-flow spectroscopy in MeCN with Zn(DMU)6(C104)2 indicate that the rate constant for reaction 9 is greater than los M-l s-l. With slow scan rates and in the absence of “filming” of the electrode by Zn02, the overall process approaches the net reaction Zn2+ + O2
+ 2e-
-
Zn02
(11)
with an O2 peak height that is equivalent to a two-electron process. However, in MeCN, DMF, and Pyr there is a strong tendency for Zn02 to precipitate on the electrode surface and
1724
ANALYTICAL CHEMISTRY, VOL. 54, NO. 11, SEPTEMBER 1982
passivate it for the O2reduction process (see Figure 4). Similar processes and problems undoubtedly occur with the other metal ions of Table IV that promote the formation of insoluble or reactive metal peroxides. In the case of Mn(I1) or Co(I1) the initial Lewis-acid promoted disproportionation of 0; is followed by further degradation steps to achieve an overall four-electron-per-02 process. For example
similar to reactions 19 and 20 have been observed with the electrochemical reduction of H202at mercury electrodes in DMF and MeCN (28). The almost reversible cyclic voltammogram for O2 at a glassy carbon electrode in MeCN indicates that specific reactions occur between metal electrodes and 0; in a poorly anionic solvating medium such as MeCN. The present results illustrate the significant effects of media and electrode material upon the reduction potentials and voltammetric peak currents for dissolved oxygen. Hence, the HzO Mn"(Oz2-) Mn" use of these parameters for the identification and quantitative determination of O2 requires substantial knowledge of the [(OH)Mn111-O-Mn111(OH)]2+ (12) media and appropriate calibration of the electrochemical [(OH)Mn111-O-Mn111(OH)]2+ 202- H20 electrode system. Some of the interferences (e.g., CH,Cl, 2Mn" 2 0 2 40H- (13) HCC13,and CC1,) may have substantial permeability through the membranes of the polarographic oxygen-membrane senThese and related reactions lead to an overall reaction for the sors (Clark electrode), which will cause substantial positive Mn"-catalyzed reduction of O2 errors. However, these problems can be turned to an advantage. The enhanced reduction currents for the reduction O2 4Mn" 2 H 2 0 4e4MnI1(OH)+ (14) of 0 2 (at 1 atm) in the presence of alkyl halides, esters, and All of these metal-oxygen intermediates are expected to acyl halides can be used for their specific determination (with be reactive toward organic substrates and electrode surfaces. appropriate calibration). Hence, the presence of metal cations enhances the electron stoichiometry for the reduction of 02,but frequently passivates LITERATURE CITED the electrode surface. (1) Maricle, D. L.; Hodgson, W. G. Anal. Chem. 1965, 3 7 , 1562. (2) Peover, M. E.; White, B. S. J. Chem. Soc., Chem. Commun. 1965, Thus, in the case of the formation of Zn02 on the surface 183. of a platinum electrode, via reactions 9 and 10, a likely process (3) Peover, M. E.; Whlte, B. S. Nectrochim. Acta 1966, 1 1 , 1061. is a metathesis reaction (4) Sawyer, D. T.; Roberts, J. L., Jr. J. Electroanal. Chem. 1966, 12, 90.
-
+
+
+
+
+
+ PtO
-
Zn(0,)
-
+ +
+
-
Pt(02)+ ZnO
(15) Likewise, a similar process may occur between the electrode surface and the H202that is produced via reactions 2 and 7
+ Hz02
-
+
Pt(O2) H2O (16) In the case of Mn(I1) ions, the presence of H202from the proton induced disproportionation of 02-(reactions 2 or 7) can lead to the production of .OHradicals PtO
Mn"
Pt
-
+ H202
+ .OH
-
+
Mn111(OH)2+ .OH
Pt(OH)*
*OH
PtO
+ H2O
(17) (18)
Because the weakest solvation of 02-occurs in MeCN (of the solvents that have been considered in the present study), superoxide should exhibit its maximum reactivity in this solvent. The peak separations for the 02/02cyclic voltammograms with metal electrodes strongly support this proposition (Figure 5 and Table I). To account for the extreme peak separation for the O2couple at platinum in MeCN, a reaction sequence that involves the formation of platinum peroxides is proposed PtO(Pt)
Ha0 + 202- -20H-2Pt(0z2-)
2Pt0
-+
.
2Pt (OH)-0-0-Pt (OH) (19) Pt(0H)-0-0-Pt(0H)
+ 20H2Pt0
O2
Johnson, E. L.; Pool, K. H.; Hamm, R. E. Anal. Chem. 1967, 39, 668. Sawyer, D. T.; Interrante, L. V. J. Nectroanal. Chem. 1961, 2 , 310. Sawyer, D. T.; Seo, E. T. Inorg. Chem. 1977, 16, 499. Bauer, D.; Beck, J.-P. J. Electroanal. Chem. Interfacial Nectrochem. 1872, 40, 233. (9) Wllshire, J.; Sawyer, D. T. Acc. Chem. Res. 1979, 72, 105. (10) Sawyer, D. T.; Nannl, E. J., Jr. "Oxygen and Oxy-Radicals in Chemistry and Blology"; Rodgers, M. A. J., Powers, E. L., Eds.; Academic Press: New York, 1981; pp 15-44. (11) Goolsby, A. D.; Sawyer, D. T. Anal. Chem. 1967, 3 4 , 411. (12) Sawyer, D. T.; Roberts, J. L., Jr. "Experimental Electrochemistry for Chemists"; Wiley-Intersclence: New York, 1974; p 46. (13) Ilan, Y. A.; Meisei, D.; Czapski, G. I s r . J. Chem. 1974, 72, 891. (14) Dietz, K.; Forno, A. E. J.: Carcombe, B. E.; Peover. M. D. J. Chem. SOC.B 1970, 816. (15) Merritt, M. V.; Sawyer, D. T. J. Org. Chem. 1970, 3 5 , 2157. (16) Sari Filllpo, J., Jr.; Chern, C.4.; Valentine, J. S. J. Org. Chem. 1975, 40. 1678 (17) Johnson,R. A.; Nldy, E. G. J. Org. Chem. 1975, 40, 1680. (18) San Fililpo, J., Jr.; Romano, L. J.; Chern, C.-I.; Valentine, J. S.J. Org. Chem . 1976, 4 1, 586. (19) Magno, F.; Bonetempeili, G. J. Nectroanal. Chem. 1976, 6 8 , 337. (20) San Fillipo, J., Jr.; Chern, C.4.; Valentine, J. S . J. Org. Chem. 1976, 41, 1077. (21) Johnson, R. A. Tetrahedron Lett. 1977, 10, 331. (22) Gibian, M. J.; Sawyer, D. T.; Ungermann, T.; Tangpoonpholvivat, R.; Morrison, M. M. J. Am. Chem. SOC. 1979, 701, 640. (23) Roberts, J. L., Jr.; Sawyer, D. T. J. Am. Chem. SOC.1961, 103, 712. (24) Yamadagni, R.; Payzant, J. D.; Kebarle, P. Can. J. Chem. 1973, 5 7 , 2507. (25) Chin, D.-H.; Chiericato, G., Jr.; Nannl, E. J., Jr.: Sawyer, D. T. J. Am. Chem. SOC. 1982, 104, 1298. (26) Barrette, W. C., Jr.; Johnson, H. W., Jr.; Sawyer, D. T. submitted for publicatlon In J . Am. Chem. SOC (27) Po, H.; Sawyer, D. T., unpublished data, 1981. (28) Morrison, M. M.; Roberts, J. L., Jr.; Sawyer, D. T. Inorg. Chem. 1979, 18. 1971. (5) (6) (7) (8)
+ 2 H 2 0 + 2e-
(20)
The latter process would be expected to occur at a more positive potential than that for the oxidation of 0,. Processes
RECEIVED for review January 11,1982. Accepted June 7,1982. This work was supported by the National Science Foundation under Grant No. CHE-7922040.
