Electrochemistry of Proton-Coupled Redox Reactions Role of the Electrode Surface H. Holden Thorp North Carolina State University, Raleigh, North Carolina 27695-8204
The cyclic voltammetry in aqueous solutions of pH-dependent redox reactions is sometimes strongly dependent on the surface properties of the working electrode. In this article, particularly striking examples of these effects are presented. Cyclic voltammograms of proton-coupled redox reactions measured at activated glassy carbon, edge-oriented pyro1yt:c graphite, and tin-doped indium oxide are discussed. These voltammograms show marked improvement in reversibilitv relative to voltammoerams measured with bare-metal & uuactivated glassy-carbon working electrodes.This imomvement is manifested either through increased current response, reduced peak-to-peak splitting, or both.
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The Importance of pH-Dependent Redox Reactions A resurgence of interest in pH-dependent redox reactions has resulted from their potential importance in catalysis, molecular electronics, and biologic~lsystems. In catalysis, metal 0x0 complexes t h a t can be reduced by combinations of electrons and protons are attractive in developing catalysts that perform oxidations by hydrogen atom and hydride transfer ( I ) . In molecular electronics, protonation potentially allows redox sit,es to serve a s "switches" by trapping electrons in a particular redox site (2). In biological systems, m e t a l l ~ e n z ~ m ewhose s, function involves t h e oxidation or reduction of a metal ion, may control the kinetics and thermodynamics of the redox step by mediating acid-base chemistry of a coordinated ligand (3). This concept is particularly appealing in light of the importance of acid-base catalysis in enzyme mechanisms. Also pH-dependent redox reactions have been considered in models for metalloenzymatic reactions that have direct proton involvement, such a s the oxidation of water (4).
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The Study of Proton-Coupled Processes Recently, the term ''proFigure 1. Cyclic voltammograms of catechol(8mM) in 0.1 M H2S04 ton-coupled" has been asat 100 mV1s. Worm electrodes cribed to these redox oroare polished glassy carbon acti- cesses (13).~t has beeLunvated by oxidation at +1.8 V in 0.1 clear whether this term M H,SO, for (a) o min, (b) 2 min, to concerted (c) 12 min, and (d) 28 min. (Reorinted with oermission from Cab one-proton, one-electron aniss (5~. cbpyright 1985 by the (i.e., hydrogen atom) transAmerican Chemical Society.) fer or to net one-proton,
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one-electron transfers, in which the proton may kinetically precede the electron or vice versa. In this article, the term "proton-coupled" is used only in the thermodynamic sense, that is, pH-dependent redox potentials are observed. This pH dependence follows the Nernst equation, giving a change of 59 mV for every pH unit for a one-proton, oneelectron couple. An attractive way to study proton-coupled redox reactions is bv cvclic voltammetrv. However. these studies are often cokpl&atedby slow electrode kinetics resultingfrom slueeish proton transfer. as in the two-electron oxidation of
The proton requirement can lead t o highly irreversible voltammograms a t electrodes made of metal or polished glassy carbon, as shown in Figure la. Activation of the carbon-electrode surface, however, permits the measurement of well-behaved reversible voltammomams. as shown in Figure Id. In this article, I will discuss electrode surfaces a t which certain proton-coupled redox reactions may be studied wlth greater facility. In the largely emipirical discuss~onI wll not seek toexplam theongm and fundamental aspens of these effects. Rather I wlll orovlde a set of tools with which chemists interested in these reactions can attempt to obtain useful electrochemical information.
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Activated Glassy Carbon Activation of freshly polished glassy-carbon electrodes by electrochemical oxidation a t +1.8 V in 0.1 M HzSOa greatly increases the reversibility of cyclic voltammograms of catechol(5). At the polished electrode, broad peaks and very large peak-to-peak splitting (AEJ are observed. After 28 min of activation a t 1.8 V, a reversible two-electron wave is evident (Fig. Id). From X-ray photoelectron spectroscopy, the origin of the effect appears to be the formation of oxidized-carbon functionalities (i.e., aldehydes, ketones, and carboxylic acids) on the surface. This formation catalyzes the proton transfer. In fact, the behavior at intermediate periods of activation (Figs. l b and lc) is consistent with catalysis of the initial oxidation step shown in eq 2, while further acti-
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to basal-plane pyrolytic graphite (6). Due to this difference in surface properties, EOPG can be used effectively for direct electrochemistry of redox proteins. The distinction between edge-oriented and basal-plane graphite seemed similar to that between activated and unactivated glassy carbon. Thus, we measured the cyclic voltammetry a t EOPG of another proton-coupled reaction (eq 6) that we were studying in reference to the oxidation of water by photosystem 11.
'fie method is also helpful fbr metal com lexes. Figure2 shows cyclic v o l t a m m o ~ a m of s l(NH3)rRu',~H2,12.at untreated and activated glassy-carbon electrodes ( 5 1 .The wave a t -0.3 V corresponds the couple given in eq 4 and exhibits quasi-reversible kinetics even a t the unactivated electrode (Fig. 2a).
Cyclic voltammograms of [(bpy)~Mn(O)~Mn03py)~l~+ at unactivated glassy-carbon and bare-metal electrodes are characterized bv lalee AE. values similar to that seen for eq 5 in ~ i ~ u r e 2 a ~ G i v a t i o nof the electrode by the +1.8N).2-V cycle method ~roduceswclic voltammoerams that exhibit quasi-reversible kinetic;; analogous obiervations are made a t EOPG (Fig. 3) (7).Detailed analysis of the temperature dependence of the voltammetry suggests that the activation barrier to proton-coupled reduction is due largely to the free energy ~ ? ~ r o t o n a t i bofnthe 0x0 species (8).
I n wntrast, the oxidation wave, corresponding to eq 5, is characterized by extremely large (610 mV a t v = 100 mV/s) peak-to-peak splitting a t the unactivated electrode. Activation (Fig. 2b) results in a decrease in AED to 190 mV. For cationic metal complexes, better results are ohtnined using activation cycles in 0.1 M H2S0, for 30 s at -1.8 V followed by 15 s ut 4 . 2 V instead of continuoui oxidative electrolgsis. Activation for about 3 such cycles appears to give the best results. This effect may be due to the localized pockets of negative charge that were created during reductive electrolysis. They may encourage more favorable interaction of the cationic complexes with the catalytic surface sites. Edge-Plane Pyrolytic Graphite The surface of edge-oriented pyrolytic graphite (EOPG) has been shown to exhibit a high oxygen content relative
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Tin-Do~edIndium Oxide Metal oxide electrodes have been used widely i n spectroelectrochemistry and in the electrochemistry of irr&ersibly adsorbed prbteins (9).At neutral pH, there is a high density of free oxide functionality on the surface, which suggested to us that these surfaces might mediate proton-coupled electron transfer. We found that the complex [Ru"(tpy)(bpy)OH2I2+exhibits two one-electron, oneproton oxidations a s shown in eqs 7 and 8 with E , ( N A I I ) = 0.62 V and Em(III/II) = 0.49 V (5)a t pH 7.
The cyclic voltammogram at tin-doped indium oxide of [ R u l ' l t ~ v l ~ h ~ v ~ Oat H ,OH I ' ~4.5 shows extremelv Door resolution bf the two adjacent couples (eqs 7 and 8)'d;e to low current response for the NAII wave (Fiz. 4a). At this DH. moht of thebxide sites on the electrode s'kface are proion: ated (101.At DII 7, however, the oxide sites are de~romnated, and the cyclic voltammogram shows two well-resolved, quasi-reversible waves (Fig. 4b). Thus, it appears that in-
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Figure 2. Cyclic voltammograms of [(NH~),RU(OH)]~+ in H 2 0 (solid line) and D 2 0 (dashed line) at (a)untreated and (b) activated glassycarbon electrodes (200 mVls). The activation was performed by two cycles of oxidation at +I .8 V for 30 s followed by reduction for 15 s a t -0.2 V in 0.1 M H2S0,. (Reprinted with permission from Cabaniss (5). Copyright 1985 by the American Chemical Society.)
Figure 3. Cyclicvoltammogram of [(bpy),~n(~)2~n(bpy]2]3+ in 0.1 M ohosohate, 0.9 M KNO, solution ioH 4.5) at an EOPG workina electrode (scan rate: lo mvis).
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time (11-13). In spite of much effort in this area, the mechanistic origin of these electrode effects remains unclear. Empirically, it appears in the cases discussed here that free~oxidcfunetionillity on the electrode surface is required for quasl-reversible electrochemistry. Thus. 11 is t e m ~ t i n g to propose that the improved electroehemistry r e s u l t s f r o ~ catalysis of the proton transfer by the oxygen lone pairs on the surfaces of activated glassy carbon, edge-oriented pyrolytic graphite, and deprotonated indium oxide. However, ~t 1s Important to punt out that many protoncoupled redox reactlond do not exhibit the affects shown in Figures 1 4 (14). I n fact, a number of proton-mupled redox reactions, including eq 4, exhibit reversible electrochemistry a t many electrodesurfaces, even a t bare metal. Thus, the involvement of a n acid-base reaction in a redox event does not necessitate the unusual electrode kinetics discussed here. In addition, a number of simple one-electron redox reactions, such a s the [Fe(CN)612Js couple, exhibit more reversible electrochemistry a t activated electrodes (5, 11-13). Although the mechanism is unclear, the fact remains that the use of the electrodes described here can a t l-v ~ ~a e~ improve voltammetric data on proton-coupled redox reactions. The use of these electrodes does not require costly equipment or special procedures, and instructions for their use can be found in the literature cited. Ho~efullv. ",this discussion clearly points out that using cyclic voltammetry to study a particular redox c o u ~ l eshould not be dismissed solel; on-the observation of poor electrode kinetics a t typical electrodes. If it is possible that a n acid-base reaction is coupled to the redox event, the electrodes discussed here may permit the obsewation of cyclic voltammetm that is more~easilyinterpreted. ~
Figure 4. Cyclic voltammograms of [ ~ u " ( t p y ) ( b p y ) ~tin-doped ~~]~at indium oxide in (a)0.1 M phosphate. pH 4.5. 20 mVIs; and (b) 0.05 M phosphate, pH 7, 25 mV/s. creasing the pH permits catalysis of the proton-coupled reactions by the deprotonated oxide functionalities. I t is intriguin that a similar transition in the cyclic voltammetry of [ ~ u K ' ( t p y ) ( b p ~ ) ~oecurs H ~ l ~upon + activating a glassy-carbon electrode (5).Thus, increasing the availablity of free "oxidelike" functionalities on the electrode surface--either by activation of the glassy-carbon electrode or deprotonation of the tin-doped indium oxide electrode-has quahtatively the same eflect on the cyclic voltammetry ol'lR~"~tpy~~b~.~,OH~I~'. Conclusions The improvement of electrode kinetics by pretreatment of carbon electrodes has been under investigation for some
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Acknowledgment I would like to thank many people: the donors of the Petroleum Research Fund administered by the American Chemical Society for financial support; the Camille and Henry Dreyfus Foundation for a New Faculty Award; the National Science Foundation for a Presidential Young Investigator Award; and E. F. Bowden, J. E. Sarneski, and J. C. Brewer for helpful discussions. Literature Cited 1. Mcyer,T. J. J.El&ckam Sac. 1984,131,221C. 2. Blaho, J. K; Coldsby, K A. J.Am Chem Soe 1890, 112,6132. 3. Thorp, H. H.: Sameski, J. E.; Bmdvig, G. W: Crabtree, R. H. J Am Chon. Soc
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4. Bmdvig, G. W ; Crabtree, R. H.Prog. 1norg Chem 1989,37,99. 5. Cabmiss, G. E.; Diamantis, A. A,: Murphy, W R., Jr;Linfon, R. W ; Meyer. T J. J , Am. Ckam Soc 1985,107,1845. 6. Armstmne.FA.:Hill. l J R R 2 1 M 7. . . . H.A. 0.: Walton.N.J.AecChem. Re% -~ 7. Thorp, H. H.; Bmdvig, G, W.: Bowden, E. F J EhelmmI. Ckan. 1890,290,293, 8. KalsbecL. W . A,: Thorp,H. H.; Bmdvig, 0.W., submitfed for publication, 9. Willit, J.L.:Bowden,E. F. J. Phys Cham. 1990,94,8241. 10. Tret'yakov, N. E.; Filimonou,V N. Kin& Catol 1972.13.735. 11. Bowling, R. J.; Packard, R. T.:MeCreery, R. L . J Am. Chem Soc 1989,111,1217. 12. Evans. J. F,h w a n a , TAnol. Chem 1977,49,1632. 13. Kepley, L.J.: Bard. A. J.Anol. Ckam. 1988,60,1159. 14. N a e ~ l iR.: , Redepnning, J.;Anron, F C. J. Phys Chem. 1986.90,6227. ~
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