Potential-, pH-, and Isotope-Dependence of Proton-Coupled Electron

Sep 21, 2006 - The Tafel plots and plots of the transfer coefficient vs overpotential are asymmetrical at all pHs. These results are interpreted in te...
0 downloads 4 Views 155KB Size
Download PDF
Langmuir 2006, 22, 10643-10651

10643

Potential-, pH-, and Isotope-Dependence of Proton-Coupled Electron Transfer of an Osmium Aquo Complex Attached to an Electrode† Nicholas Madhiri and Harry O. Finklea* C. Eugene Bennett Department of Chemistry, West Virginia UniVersity, Morgantown, West Virginia 26506 ReceiVed April 24, 2006. In Final Form: August 25, 2006 An osmium complex, [OsII(bpy)2(4-aminomethylpyridine)(H2O)]2+, is attached to a mixed self-assembled monolayer on a gold electrode. The complex exhibits 1-electron, 1-proton redox chemistry (OsIII(OH)/OsII(H2O)) at pHs and potentials that are experimentally accessible with gold electrodes in aqueous electrolytes. The thermodynamic behavior and kinetic behavior of the system are investigated as a function of pH in both H2O and D2O. The two formal potentials and two pKa values are relatively constant for two chain lengths in H2O and in D2O. The standard rate constants at all pHs are strongly and uniformly affected by chain length, indicating that electronic coupling is the dominant factor controlling the rate of electron transfer. In both H2O and D2O, the standard rate constant is weakly dependent on the pH, exhibiting a minimum value midway between the pKa values. The kinetic isotope effect is small; standard rate constants decrease by roughly a factor of 2 in D2O over a wide range of pHs, but not at the more acidic pHs. The Tafel plots and plots of the transfer coefficient vs overpotential are asymmetrical at all pHs. These results are interpreted in terms of a larger reorganization energy for the OsII species and a smaller reorganization energy for the OsIII species. The OsIII reorganization energy is constant at all pHs in both H2O and D2O. The pH dependence of the OsII reorganization energy accounts for some or all of the pH dependence of the standard rate constant in H2O and D2O. The data deviate substantially from predictions of the stepwise proton-coupled electron-transfer mechanism. The observation of a kinetic isotope effect supports the concerted mechanism.

Introduction Proton-coupled electron transfer (PCET) reactions play a critical role in a variety of biological and chemical processes.1 Recent theoretical1-12 and experimental studies13-17 have focused on homogeneous systems. These studies often discuss the relative importance of the stepwise vs the concerted mechanisms for proton and electron transfer. In the stepwise mechanism, the electron and proton are transferred in separate steps, and a chemical intermediate exists between the oxidized deprotonated form and the reduced protonated form. In the concerted mechanism, the electron and proton are transferred in the same rate-determining step. Most of the recent studies emphasize the †

Part of the Electrochemistry special issue. * To whom correspondence should be addressed. E-mail: Harry.Finklea@ mail.wvu.edu. (1) Cukier, R. I.; Nocera, D. G. Annu. ReV. Phys. Chem. 1998, 49, 337-369. (2) Chang, C. J.; Chang, M. C. Y.; Damrauer, N. H.; Nocera, D. G. Biochim. Biophys. Acta 2004, 1655, 13-28. (3) Mayer, J. M.; Rhile, I. J. Biochim. Biophys. Acta 2004, 1655, 51-58. (4) Mayer, J. M. Annu. ReV. Phys. Chem. 2004, 55, 363-390. (5) Soudachov, A. V.; Hammes-Schiffer, S. J. Chem. Phys. 2000, 113, 23852396. (6) Decornez, H.; Hammes-Schiffer, S. J. Phys. Chem. A 2000, 104, 93709384. (7) Iordanova, N.; Decornez, H.; Hammes-Schiffer, S. J. Am. Chem. Soc. 2001, 123, 3723-3733. (8) Hammes-Schiffer, S. Acc. Chem. Res. 2001, 34, 273-281. (9) Iordanova, N.; Hammes-Schiffer, S. J. Am. Chem. Soc. 2002, 124, 48484856. (10) Soudackov, A.; Hatcher, E.; Hammes-Schiffer, S. J. Chem. Phys. 2005, 122, 014505. (11) Hatcher, E.; Soudackov, A.; Hammes-Schiffer, S. Chem. Phys. 2005, 319, 93-100. (12) Hatcher, E.; Soudackov, A.; Hammes-Schiffer, S. J. Phys. Chem. B 2005, 109, 18565-18574. (13) Roth, J. P.; Lovell, S.; Mayer, J. M. J. Am. Chem. Soc. 2000, 122, 548698. (14) Yoder, J. C.; Roth, J. P.; Gussenhoven, E. M.; Larsen, A. S.; Mayer, J. M. J. Am. Chem. Soc. 2003, 125, 2629-2640. (15) Soper, J. D.; Mayer, J. M. J. Am. Chem. Soc. 2003, 125, 12217-12229. (16) Rhile, I. J.; Mayer, J. M. J. Am. Chem. Soc. 2004, 126, 12718-12719. (17) Lebeau, E. L.; Binstead, R. A.; Meyer, T. J. J. Am. Chem. Soc. 2001, 123, 10535-10544.

importance of the concerted mechanism. In that mechanism, large kinetic isotope effects are possible if the proton or deuteron tunnels over a large distance.18-20 More recently, theoretical analyses of PCET for heterogeneous electron transfer have been published21,22 and applied to two systems.23,24 The analysis by Constentin and co-workers leads to the following overall expression for the cathodic rate constant kc22

∫

kc ) Z(RT/4πλ)0.5 D(λ,η,ζ) f(ζ) dζ

(1)

D(λ,η,ζ) ) exp{-(RT/4λ)[λ/RT + Fη/RT - ζ]2}

(2)

f(ζ) ) (1 + exp(ζ)) ζ ) ( - F)/RT

-1

(3) (4)

where λ is the reorganization energy, η is the overpotential (E - E°),  is the energy, and F is the Fermi energy of the metal electrode. The preexponential term Z includes the density of states in the metal and the electronic coupling between the electrode and the redox molecule. Equation 1 has the same form as the equation used for simple electron transfer.25 D(λ,η,ζ) can be described as the density of acceptor states associated with the oxidized form of the redox couple and f(ζ) is the Fermi function (18) Huynk, M. V.; White, P. S.; Meyer, T. J. Angew. Chem., Int. Ed. 2000, 39, 4101-4410. (19) Huynk, M. H. V.; Meyer, T. J. Angew. Chem., Int. Ed. 2002, 41, 13951398. (20) Huynh, M. H. V.; Meyer, T. J. Proc. Natl. Acad. Sci. 2004, 101, 1313813141. (21) Grimminger, J.; Bartenschlager, S.; Schmickler, W. Chem. Phys. Lett. 2005, 416, 316-320. (22) Costentin, C.; Robert, M.; Saveant, J.-M. J. Electroanal. Chem. 2006, 558, 197-206. (23) Constentin, C.; Evans, D. H.; Robert, M.; Saveant, J.-M.; Singh, P. S. J. Am. Chem. Soc. 2005, 127, 12490-12491. (24) Costentin, C.; Robert, M.; Saveant, J.-M. J. Am. Chem. Soc. 2006, 128, 4552-3. (25) Finklea, H. O. In Electroanalytical Chemistry; Bard, A. J., Rubinstein, I., Eds.; Marcel Dekker: New York, 1996; Vol. 19, pp 109-335.

10.1021/la061103j CCC: $33.50 © 2006 American Chemical Society Published on Web 09/21/2006

10644 Langmuir, Vol. 22, No. 25, 2006

Madhiri and Finklea

Figure 2. Square scheme for a 1e1H redox system, represented by the osmium aquo complex. Table 1. Thermodynamic Parameters for the [OsII/III(bpy)2(4-Amp)(L)] (L ) OH- or H2O) Complex

Figure 1. Self-assembled monolayer of [Os(bpy)2(4-aminomethylpyridine)(H2O)]2+ bonded to HS(CH2)15COOH on gold. The diluent thiol is HS(CH2)16OH.

associated with the donor states of the metal. The reorganization energy is the sum of three components, the inner reorganization energy (λi) and solvent reorganization energies associated with electron (λET) and proton transfer (λPT). Methods are given for the calculation of the three components of the reorganization energy and the preexponential factor Z. The kinetic isotope effect (ratio of standard rate constant in H2O to the standard rate constant in D2O) is obtained from the ratio of the Z factors. This paper is part of an ongoing study of heterogeneous electron-transfer reactions with coupled proton transfer.26-29 The study takes advantage of two beneficial effects of self-assembled monolayers as spacers between the metal and the redox center. Tunneling across the monolayer slows the electron transfer rates to experimentally accessible ranges, and the low dielectric constant of the monolayer minimizes double layer effects in the kinetic measurements.25 Consequently, rates of electron transfer can be measured over a wide range of overpotentials for both the oxidation and the reduction steps. The potential dependence of the rate constants is directly interpretable in terms of Marcus theory reorganization energies (λ). With the choice of the right model compound, [OsII/III(bpy)2(4-AMP)(L)] (L ) OH- or H2O; 4-AMP ) 4-aminomethylpyridine) (Figure 1), kinetic parameters can be measured over a range of pHs that encompass all four species of a 1-electron,1proton redox couple (Figure 2). The thermodynamic parameters (two formal potentials and two acid dissociation constants) for this complex fall in ranges that are conveniently accessible for gold electrodes covered with self-assembled monolayers (Table 1). In the previous paper,29 the thermodynamic and kinetic parameters of the redox molecule were measured as a function of potential and pH when the molecule was coupled to a thinner mixed monolayer containing HS(CH2)15COOH and HS(CH2)12OH on gold electrodes. The kinetic data exhibited substantial deviations from the predictions of the stepwise PCET mechanism. Specifically, the standard rate constant ks and the transfer coefficient measured at the pH-dependent formal potential R(0) (26) Finklea, H. O.; Haddox, R. Phys. Chem. Chem. Phys. 2001, 3, 34313436. (27) Finklea, H. O. J. Phys. Chem. B 2001, 105, 8685-8693. (28) Haddox, R. M.; Finklea, H. O. J. Electroanal. Chem. 2003, 550-551, 351-358. (29) Haddox, R. M.; Finklea, H. O. J. Phys. Chem. B 2004, 108, 1694-1700.

solvent

SAMb

pKa1

pKa2

E01 c

E02 c

ref

H2O H 2O H 2O D 2O D 2O d

none thinner thicker thicker thicker

2.15 2.4 2.05 2.44 2.67

9.74 9.3 9.70 10.21 9.90

+0.32 +0.34 +0.41 +0.42

-0.13 -0.07 -0.04 -0.04

this work 29 this work this work

a Thermodynamic parameters are obtained by fitting a plot of the formal potential vs pH to the 1e1H model. See Figure 7 and the Supporting Information. b Thinner self-assembled monolayers contained HS(CH2)15COOH and HS(CH2)12OH, thicker SAMs contained HS(CH2)15COOH and HS(CH2)16OH. Values measured on a bare gold electrode for the complex dissolved in the electrolyte are shown on the first line. c Formal potentials with respect to the Ag/AgCl reference electrode. d The pKa values measured in D2O are corrected to the pKa values in H2O using the linear correlation reported by Krezel and Bal.34

showed a much weaker pH dependence than predicted. The results led to the tentative hypothesis that the mechanism was concerted rather than stepwise. However, the data were not explicitly interpreted in terms of the concerted mechanism. In this paper, the electron transfer kinetics of the osmium 1e1H redox couple is examined using a slightly thicker SAM containing HS(CH2)15COOH and HS(CH2)16OH (Figure 1). For a simple electron-transfer system, this increase in thickness using a longer diluent thiol decreases the electronic coupling between the redox center and the electrode and slows the standard rate constant significantly.30 In the osmium system, the slower standard rate constant for the thicker SAM permits the measurement of rate constants to larger overpotentials and, consequently, provides a better opportunity to examine the reorganization energies of the OsIII and OsII forms. The data analysis starts with eq 1. That equation, when translated into the standard rate constant expression (kc ) ks exp(-Rfη), where f ) F/RT and η ) E - E°) yields a potential dependent transfer coefficient R. A plot of the transfer coefficient R vs the overpotential η provides a more facile route to the measurement of the reorganization energies than a Tafel plot. For the case of equal reorganization energies for the oxidized and reduced form, the plot is nearly linear over a range of positive and negative overpotentials, and the slope of the plot can be translated into the reorganization energy. The data obtained in this study indicate that the reorganization energies of the oxidized and reduced forms are different.31 The effect of unequal λ values on Tafel plots and R vs η plots is given in the Supporting Information. For the case of small differences between large reorganization energies, the R vs η plot is again nearly linear but with different slopes at positive and negative overpotentials. The slope of the plot at positive potentials can be converted to the reorganization energy for the oxidized form Os(III). Similarly, (30) Finklea, H. O.; Liu, L.; Ravenscroft, M. S.; Punturi, S. J. Phys. Chem. 1996, 100, 18852-18858. (31) Hupp, J. T.; Weaver, M. J. J. Phys. Chem. 1984, 88, 6128-6135.

Proton-Coupled Electron Transfer at a SAM

Figure 3. Cyclic voltammetry of [Os(bpy)2(4-AMP)(H2O)]2+ at a bare gold bead electrode in 0.5 M K2SO4 + 0.1 M Britton-Robinson buffer. Scan rate 10 mV/s. The lower right CV was recorded at pH 4.4. The upper left CV (offset up by 0.5 µA for clarity) was recorded at pH 12.4. Each CV shows the OsII/III and OsIII/IV redox couples.

the slope of the negative branch of the plot yields the reorganization energy of the reduced form Os(II). Also, R(0) is determined by the two reorganization energies. In particular, when λOx is smaller than λRed, R(0) is less than 0.5. Because the kinetic isotope effect is often used as a diagnostic for the concerted PCET mechanism,22 the kinetic measurements have been performed in both H2O and D2O electrolytes. The following questions will be addressed: (a) How are the thermodynamic parameters affected by the chain length of the SAM in H2O? (b) How are the thermodynamic parameters affected by replacing the proton with the deuteron? (c) How are the kinetic parameters affected by the chain length of the SAM and the pH in H2O? (d) How are the kinetic parameters affected by replacing the proton with the deuteron? (e) Can the data be rationally explained by a concerted PCET mechanism? Experimental Section All chemicals were purchased and used as received. Synthesis and Characterization of [OsII(bpy)2(4-AMP)(H2O)](PF6)2. The syntheses29 are described in more detail in the Supporting Information. (NH4)2OsCl6 and 2,2′-bipyridine (bpy) were heated to reflux in deoxygenated DMF to produce Os(bpy)2Cl2. Os(bpy)2CO3 was synthesized by refluxing Os(bpy)2Cl2 in a concentrated aqueous Na2CO3 solution. Dissolution of Os(bpy)2CO3 in weakly acidic aqueous solution rapidly converted it to Os(bpy)2(H2O)22+. Addition of 4-aminomethylpyridine and refluxing yielded the desired product, which was isolated as the PF6- salt. At each stage, the product identity was demonstrated by its electrochemical behavior in an appropriate electrolyte and by UV/vis spectra. NMR data (Varian Inova 600 MHz) are included in the Supporting Information. Figure 3 shows the cyclic voltammograms of the product in aqueous buffers at two pHs. Both the OsIII-OH/OsII-OH2 and OsIV ) O/OsIII-OH redox couples are visible. The two waves shift with pH with a slope of -59 mV/pH. Table 1 contains the pKa values and formal potentials for the OsII/III redox couple. At pHs above 12, the two waves start to merge to form a 2-electron, 2-proton OsII/IV redox wave. This behavior limits kinetic experiments on the OsII/III wave to pHs below 12. Preparation of the Electrodes, SAM Deposition, and the Coupling Reaction. 16-Mercaptohexadecanoic acid, HS(CH2)15COOH (5 mg, 17.3 µmol), was dissolved in 10 mL of absolute ethanol. 16-Mercapto-1-hexadecanol, HS(CH2)16OH (5 mg, 18.2 µmol), was also dissolved in 10 mL of absolute ethanol. A monolayer deposition solution was made by mixing 190 µL of the HS(CH2)15COOH thiol with 10 µL of the HS(CH2)16OH thiol in a vial containing 2% (v/v) of trifluoroacetic acid in 10 mL of absolute ethanol. The final concentration of HS(CH2)15COOH was 33 µM while that of HS(CH2)16OH was 1.8 µM.

Langmuir, Vol. 22, No. 25, 2006 10645 Polycrystalline gold bead electrodes were fabricated by cutting a 0.5 mm gold wire (99.999%) into 3 cm lengths. One end of the wire was melted in a Bunsen burner flame to form the bead. The average bead diameter was 1 mm. No insulation was used to define the area of the gold in contact with the electrolyte. When the bead and approximately 0.5 mm of the wire was immersed in the electrolyte, the corresponding surface area in contact with the electrolyte was about 0.03 cm2. The electrodes were prepared by chemical etching in hot “piranha” solution (3:1 volume mixture of concentrated sulfuric acid/30% hydrogen peroxide: CAUTION). They were rinsed with copious amounts of water followed by electrochemical cycling between -0.1 and 1.4 V vs Ag/AgCl in 0.1 M H2SO4. After this electrochemical step, the electrodes are rinsed again with water followed by ethanol. They were blown dry before being immersed in the thiol deposition solution overnight. The electrodes were removed from the deposition solution and washed with ethanol followed by two rinses with 10% ammonium hydroxide/ ethanol solution to remove residual trifluoroacetic acid. This rinsing procedure is believed to inhibit the formation of a second thiol layer.32 This step was followed by another rinse with 2% trifluoroacetic acid/ethanol solution to ensure protonation of the carboxylic acid terminal groups of the thiols. Electrodes were placed back in the deposition solution for several hours before coupling was done. When electrodes were ready for coupling, the same rinsing procedure was repeated, but, after the ammonium hydroxide rinse, electrodes were rinsed with water and placed in the coupling solution. This procedure resulted in monolayers of good quality as judged by low and flat charging currents in a CV. About 2-3 mg of the osmium aquo complex was dissolved in approximately 2 mL of acetonitrile to which 5 drops of water had been added. To this solution was added 10 mg of 1-(3-dimethylaminopropyl)-3-ethylcarbodiimide hydrochloride. Electrodes were placed in the coupling solution for 30 min. They were subsequently rinsed copiously with water before use. With this procedure, coverages of the osmium complex on the electrode were approximately 1 × 10-11 mol/cm2. The procedure also minimized the appearance of a persistent impurity wave in the CVs at a potential slightly negative of the OsII/III wave at lower pHs. The impurity was never identified. Electrodes were stored in water until they were ready for use. Prolonged exposure of electrode to air resulted in the monolayer being prone to degraded behavior after only a few CVs in different electrolytes. Oxidation of the thiol by traces of ozone may have caused the instability.33 Electrochemical Measurements. A three-electrode system was used in a single compartment polarographic cell (Princeton Applied Research) with a 4 mL working volume and a custom Teflon lid. In the cell, the uncompensated resistance was reduced by minimizing the distance between the reference electrode (Ag/AgCl microelectrode, Cypress) and the gold bead working electrode. A platinum flag (area 2 cm2) served as the counter electrode. All electrolyte solutions were prepared using distilled deionized water (resistivity greater than 16 MΩ cm). The electrolyte was prepared starting with 1 M H2SO4 and 0.1 M each of tri-sodium phosphate, oven-dried sodium citrate, and sodium borate (BrittonRobinson buffer), to ensure that the electrolyte was buffered over the whole pH range (1-11) investigated. About 4 mL of the stock buffer solution was added to a 20 mL vial, and the pH was adjusted using a few drops of either concentrated sulfuric acid or potassium hydroxide. The buffers were allowed to equilibrate at room temperature (22 °C) prior to use. Concentrated solutions were used to minimize uncompensated resistance in the cell and double layer effects, and to avoid significant changes in the ionic strength of the electrolyte during pH adjustment. The total ionic strength exceeded 3.0 M at all pHs. AC impedance measurements at 20-100 kHz yielded uncompensated resistances of 9 (pH 2), 10 (pH 6), and 11 Ω (pH 10) for a SAM-coated bead. Potassium hydroxide was used in preference to sodium hydroxide in order to decrease sodium error (32) Wang, H.; Chen, S.; Li, L.; Jiang, S. Langmuir 2005, 21, 2633-2636. (33) Schoenfisch, M. H.; Pemberton, J. J. Am. Chem. Soc. 1998, 120, 45024513.

10646 Langmuir, Vol. 22, No. 25, 2006

Madhiri and Finklea

Figure 5. Data analysis example. The plot of R vs η is fitted with separate linear regression lines for the positive and negative branches. The standard rate constant is adjusted so that the two linear regression lines share a common intercept R(0). These data were obtained from the CVs at pH 6.0 in Figure 4. See the text for details. Fitted parameters: ks ) 0.92 s-1, positive branch (solid line) slope ) 0.33 ( 0.05 V-1, negative branch (dashed line) slope ) 0.22 ( 0.03 V-1, R(0) ) 0.44 ( 0.02. The slopes correspond to reorganization energies of 0.64 ( 0.10 eV for OsIII and 1.00 ( 0.16 eV for OsII.

Figure 4. Overlay cyclic voltammograms in H2O buffer. Each overlay includes CVs at 0.1, 1, 10, 100, and 1000 V/s. The overlays were obtained at pHs 1.0, 2.1, 4.0, 6.0, 8.2, 10.2, and 11.0 (bottom to top). The current (µA/(V/s)) is normalized with respect to scan rate. at high pHs. The buffers were sparged using either nitrogen or argon just before use to remove dissolved oxygen. For deuterated electrolytes, the procedure was performed using D2SO4 and KOD. For all pH measurements, the combination glass pH electrode (Fisher) was calibrated in standard pH 4, 7, and 10 buffers (Fisher). All data were collected using cyclic voltammetry on a Princeton Applied Research model 283 potentiostat/galvanostat connected to a Tektronix TDS430A digital oscilloscope and a PAR Model 175 Waveform Generator. CVs were collected at scan rates of 0.1, 1, 10, 100, and 1000 V/s with typically 1000 data points per CV. Figure 4 shows overlay plots at pHs 1, 2, 4, 6, 8, 10, and 11 in aqueous buffer, and measurements were also obtained at intervening pHs. The initial potential was set at a value negative of the OsII/III formal potential, scanned to a value sufficiently positive to fully oxidize all of the osmium complex to OsIII and then to a value sufficiently negative to fully reduce all of the osmium complex to OsII, and returned to the initial potential. At the faster scan rates of 100 and 1000 V/s, the switching potentials were usually beyond the potentials associated with gold oxidation (positive limit) or monolayer desorption (negative limit, especially at the higher pHs),25 but the rapid potential excursion at these scan rates did not result in any significant degradation of the monolayer as indicated by the absence of changes in the coverage of the redox centers or changes in the charging current on subsequent CVs. Evidently, thiol desorption does not occur within the millisecond during which the electrode potential is near the negative potential limit. Data Analysis. Data analysis assumes the validity of the Marcus model with different reorganization energies for the OsII and OsIII species. All calculations were done on a spreadsheet. The anodic and cathodic currents were separately analyzed. A charging current baseline was extrapolated under the faradaic peak and subtracted to isolate the faradaic current. The peak area was integrated to obtain the total faradaic charge. The formal potential was obtained from the average of the anodic and cathodic peak potentials at 0.1 V/s; peak splittings were typically less than 100 mV at this scan rate. The kinetic data was obtained at 10, 100, and 1000 V/s. The potential and current corresponding to 50% of the faradaic

Figure 6. Tafel plot for the data in Figure 4 using the parameters listed in Figure 5. charge remaining were located. The potential was converted to overpotential with inclusion of an iR drop correction based on 10 ohms uncompensated resistance. The rate constant was obtained by dividing the faradaic current by the amount of faradaic charge remaining. This procedure yielded a single anodic and cathodic rate constant and corresponding overpotential per scan rate. See the Supporting Information for a more detailed description of this procedure. The procedure of analyzing the kinetics at one point on the oxidation or reduction wave was adopted in order to obtain the average kinetic behavior of the redox centers. In these experiments as well as in preceding experiments with different redox centers, the analysis of the data is complicated by kinetic heterogeneity.26,28,29 The redox centers exhibit a distribution of standard rate constants (and possibly of reorganization energies) rather than a uniform value. In this situation, analysis of the anodic wave or cathodic wave by comparison with a simulated linear scan voltammogram works well for a single scan rate, but does not work well for all of the scan rates at a given pH (see the Supporting Information). A single set of kinetic parameters (one standard rate constant and two reorganization energies) does not provide a consistently close fit of simulated curves to anodic or cathodic data at 10, 100, and 1000 V/s. The choice of kinetic analysis at the point at which 50% of the faradaic charge is remaining is based on simulations of kinetic heterogeneity; that point yields kinetic parameters that are close to the average at all scan rates. That point is generally located near the peak faradaic current. Consequently, uncertainties introduced by the choice of the charging current baseline are minimized. The pairs of rate constants and overpotentials were converted to a plot of the transfer coefficient R vs η using an adjustable standard

Proton-Coupled Electron Transfer at a SAM

Langmuir, Vol. 22, No. 25, 2006 10647

rate constant ks. Linear regression lines were calculated for the positive branch and negative branches of the plot. The standard rate constant was adjusted until the two linear regression lines shared a common intercept at η ) 0; see Figure 5 for an example. Figure 6 shows the corresponding Tafel plot.

Results and Discussion A. Reversible Behavior and Double Layer Effects. It is worth examining how close the Os complex electrochemistry is to ideal thermodynamic behavior in CVs recorded at 0.1 V/s. Because iR drop is negligible at this scan rate, a simple simulation procedure can be used to calculate linear scan voltammograms with matching peak areas and peak positions (see the Supporting Information). Deviations of peak half-widths from predicted values can indicate problems that might compromise the kinetic analyses. Ideal peak half-widths vary from 92-95 mV for standard rate constants of 5-7 s-1 to 100-110 mV for standard rate constants of 1-2 s-1. These ranges of standard rate constants are typical for the Os redox couples in this SAM (see below). At pHs between 1 and 4, the anodic and cathodic peak half-widths are both in the range of 115 ( 10 mV (Figure 4), somewhat higher than the ideal value. This behavior is seen for other redox centers and is usually attributed to a slight variation of formal potential (thermodynamic heterogeneity) or interactions between redox centers (possibly indicating clustering).25 More puzzling is the behavior at pHs above 4 (Figure 4). The anodic peak half-widths decrease to 105 ( 10 mV (close to the ideal half-widths) and the cathodic peak half-widths increase to 125 ( 10 mV. The results in the next section show that the higher pH range corresponds to OsIII-OH being the oxidized form and OsII-H2O or OsII-OH being the reduced form of the redox center. Nonideal behavior can result from increased partitioning into the outer surface of the SAM. It is not clear why the higher oxidation state (which presumably is more hydrophilic) would exhibit less ideal behavior due to partitioning into the SAM. Other possible sources of nonideal peak half-widths are double layer effects. Smith and White described two types of double layer effects in terms of basic electrostatic theory.34,35 For SAMs, double layer effects arise from changes of the electrostatic potential at the external surface of the SAM. If the SAM contains a redox center, then oxidation or reduction changes the charge density at the plane of electron transfer, and the electrostatic potential at that plane, φPET, changes as the applied potential scans across the formal potential of the redox center.34 The driving force for electron transfer is the difference in electrostatic potential between the metal and the plane of electron transfer (φM - φPET). The equations predict an increase in the half-width of the reversible faradaic peak because changes in the driving force are different from changes in the externally applied potential. Alternately, if the external part of the SAM contains an acidic group, such as a carboxylic acid, then the ionization of the acidic group will vary with potential as well as pH, and the electrostatic potential at the plane of acid dissociation, φPAD, will change.35 If a redox center is also present, then the variation of φPAD with pH may be evident in the dependence of the formal potential with pH. The effect of applied potential and pH on the reversible behavior of an attached redox center can be calculated given the values of a number of parameters. For the experiments reported in this paper, the following representative parameters are used: thickness of SAM 1.5 nm (a 15-methylene alkanethiol with a terminal group and an average tilt of 30 degrees from the vertical would (34) Smith, C. P.; White, H. S. Anal. Chem. 1992, 64, 2398-2405. (35) Smith, C. P.; White, H. S. Langmuir 1993, 9, 1-3.

Figure 7. Formal potential E° vs pH for [Os(bpy)2(4-AMP)(H2O)]2+ attached to the thicker SAM. The solid line is a least-squares fit to the 1e1H model. See Table 1 for the fitted parameters. The pKa values are marked on the pH axis.

be 1.8 nm thick), dielectric constant of the SAM 8 (based on observed charging currents), and an inverse Debye length of 3.3 × 109 m-1 (a value typical of a 1 M 1:1 electrolyte). For the electroactive SAM, the coverage was 1 × 10-11 mol/cm2, and the charges of the oxidized form and reduced forms were +2 and +1, respectively. These charges are correct for pHs above pKa2. At pH less than pKa1, the charges would be +3 and +2, which would result in smaller distortions from ideal behavior. At pH between the two pKa values, the charges are +2 in both oxidation states, and there would be no double layer effect at all. Given these parameters, the equations show no significant changes in the peak half-width for formal potentials between +1 and -1 V vs Epzc, the potential of zero charge. The high ionic strength and low redox coverage minimizes the variation of φPET with applied potential. Consequently, this source of peak broadening is considered to be negligible. The SAM contains a mixture of HS(CH2)16OH and HS(CH2)15COOH. Some of the carboxylic acids are left unreacted after the coupling of the Os complex. The ionization of these acid groups depends on both the pH and the applied potential. Using the same SAM thickness, SAM dielectric constant, and electrolyte inverse Debye length, and assuming a coverage of 4 × 10-10 mol/cm2 of ionizable COOH groups with pKa ) 5.0, the equations given by Smith and White35 predict a slope of φPAD with respect to applied potential of +20 mV per V at low pH (no ionized sites) and a slope of +6 mV per V at high pH (fully ionized sites). For a redox peak with an ideal half-width of 90-100 mV, the variation in φPAD would increase the half-width by 2 mV or less. Again, the high ionic strength and low coverage of redox centers minimizes the impact of double layer effects on reversible behavior. We conclude that the pattern of peak half-widths for the Os redox centers at 0.1 V/s must arise from factors other than double layer effects. Further evidence for this conclusion is given in results previously reported for the [OsII/III(bpy)2(4-AMP)(Cl)]+/2+ complex attached to a 100% HS(CH2)15COOH SAM.29 In 0.5 M K2SO4 + 0.1 M Britton-Robinson buffer, the formal potential was constant with a standard deviation of (7 mV for pHs between 2 and 9. These results indicate that the electrostatic equations of Smith and White34,35 may overestimate the magnitude of double layer effects. The potential impact of double layer effects on kinetic behavior is discussed later. B. Thermodynamic Parameters in H2O versus Chain Length. A plot of formal potential vs pH for the Os complex on the SAM can be fitted to extract the four thermodynamic constants (Figure 7). These values compare well with the solution values (Table 1) given that the pendant amine group has a pKa

10648 Langmuir, Vol. 22, No. 25, 2006

Madhiri and Finklea

near pKa2. The pKa values reported for the thinner SAM were slightly closer to each other.29 Arguments can be made for some variation in pKa values for the different SAMs. However, other data sets indicate that pKa values have an experimental uncertainty of at least (0.2. Consequently, the pKa values are, to a first approximation, independent of chain length. C. Thermodynamic Parameters in D2O versus H2O. As noted in the Experimental Section, the pD values of the electrolyte were measured using a glass pH electrode which was standardized in H2O buffers. Several papers in the literature tackle the issue of pH measurement in D2O.36-38 Krezel and Bal established that a linear correlation exists between pKa values obtained in D2O and pKa values in H2O:36

pKa(H2O) ) 0.929 pKa(D2O) + 0.41

(5)

Figure 8. log(ks) vs pH for the thicker (circles) and the thinner (diamonds) SAM. The circles are the average log(ks) at each pH for three sets of experiments; the bars are one standard deviation. The pKa values for the two data sets are marked on the pH axis.

pKa(D2O) is the value based on pD readings in D2O. Table 1 lists the thermodynamic parameters measured by a plot of formal potential vs pD in D2O electrolyte. The pKa values in D2O are somewhat higher than in H2O. Upon applying eq 5 (see the last row of Table 1), the pKa2 values measured in H2O and D2O are within the experimental uncertainty of the measurement. It is possible that the pKa1 values are also within experimental uncertainty, but more data would be needed to make that judgment. D. Kinetic Parameters in H2O versus Chain Length. As a prelude to the following discussion, it is helpful to review the method of obtaining the kinetic parameters and to consider some alternate possible causes of potential dependence of the rate constants. Each rate constant is measured at a particular potential corresponding to 50% conversion of the reactant oxidation state to the product oxidation state. This approach is believed to yield the best average kinetic behavior in a system that exhibits kinetic heterogeneity. The analysis method takes into account the effects of uncompensated resistance, which is only significant at the fastest scan rate (1000 V/s). One added factor to consider is double layer effects. The equations published by Smith and White34,35 yield an estimate of the electrostatic potential at the surface of the SAM (φPET or φPAD) relative to the electrostatic potential in the bulk of the electrolyte. Given the wide range of overpotentials that appears in the kinetic data (up to (0.8 V), it is conceivable that the magnitude of change in φPET or φPAD would affect the actual change in overpotentials. The calculations discussed in part A can be used to address this issue. As noted earlier, φPAD changes linearly with applied potential, and the corresponding derivative of φPAD with respect to the applied potential is 20 mV/V at low pH and 6 mV/V at high pH. The slope of φPET with respect to the applied potential fluctuates at potentials near the formal potential. For kinetic measurements at overpotentials greater than 0.2 V, the slope varies from 0 to +15 mV/V at all potentials within +1 or -1 V with respect to Epzc. These slopes are considered to be upper limits due to the conservative parameters used in the calculations. Their effect on the kinetic parameters would be to reduce the overpotentials by a factor of (1 - slope)η. By including the double layer correction, Tafel plots would become steeper, and the resulting reorganization energies would increase slightly. Since the slopes are either constant with applied potential (φPAD) or symmetrical with respect to the formal potential (φPET), no asymmetry would be introduced into the Tafel plots. Again, experimental data indicate that the double layer effects are negligible. The standard rate constant, Tafel slopes, and reorganization energies of an Os-Cl complex

on a HS(CH2)15COOH SAM are, within experimental uncertainty, constant from pH 1.5 to 11.29 Another factor which could impact the kinetic analysis is thermodynamic heterogeneity. In part A, the half-width of the reversible reduction wave was found to be significantly larger than the half-width of the reversible oxidation wave at pHs above 4. If the cause is a range of formal potentials for the reduction of the OsIII species, then the kinetic heterogeneity is exacerbated for the cathodic rate constants. The cathodic rate constants are translated into reorganization energies for the OsII species. Consequently, the OsII reorganization energies are likely to have greater uncertainty than the corresponding OsIII reorganization energies. Finally, we consider the validity of the assumption that the Z term in Equation 1 (the density of metal states and the electronic coupling) is independent of electrode potential over the range of potentials explored in this work. Figure 4 demonstrates that the CVs covered potentials from -1.3 to +1.3 V vs Ag/AgCl. The peak potentials (approximate markers for the overpotentials at 50% conversion) span a range of 1.7 V. There are relatively few measurements of the electronic coupling over a wide range of electrode potentials. Gosavi and Marcus calculate a density of states for gold that is nearly constant over more than 2 eV.39 Becka and Miller observed a slight increase in the tunneling coefficient with increasing potential over a potential range exceeding 1 V for blocking SAMs of two chain lengths and four solution redox couples.40 Their data were not corrected for any change in reorganization energy with respect to chain length. Similarly, Robinson and Chidsey measured the same tunneling coefficient for ferrocene on different chain length SAMs out to overpotentials of (0.6 V.41 These results support the assumption that Z is constant, but, clearly, that assumption deserves further testing. The standard rate constants vs pH for the thinner and thicker SAMs are compared in Figure 8. The difference between the standard rate constants for the two chain lengths is about a factor of 10 over the entire pH range. Both data sets exhibit a shallow minimum at a pH roughly midway between the pKa values. Neither data set exhibits the V-shaped plot predicted by the stepwise mechanism (see the Supporting Information), nor is there a clear change in the slopes of the data around the two pKa values. We conclude that the pH has at most a small effect on the electronic coupling between the Os redox center and the electrode (see below). The change in electronic coupling with respect to the increased length of the diluent thiol is consistent with the previous study

(36) Krezel, A.; Bal, J. J. Inorg. Biochem. 2004, 98, 161-166. (37) Covington, A. K.; Pabo, M.; Robinson, R. A.; Bates, R. G. Anal. Chem. 1968, 40, 700-706. (38) Baucke, F. G. K. J. Phys. Chem. B 1998, 102, 4835-4841.

(39) Gosavi, S.; Marcus, R. A. J. Phys. Chem. B 2000, 104, 2067-2072. (40) Becka, A. M.; Miller, C. J. J. Phys. Chem. 1992, 96, 2657-2668. (41) Robinson, D. B.; Chidsey, C. E. D. J. Phys. Chem. B 2002, 106, 1070610713

Proton-Coupled Electron Transfer at a SAM

Figure 9. Reorganization energies vs pH in H2O. Squares, λ for OsIII; diamonds, λ for OsII. The open symbols are the averages for three sets of experiments; bars are one standard deviation. The pKa values are marked on the pH axis.

Figure 10. R(0) vs pH in H2O. Squares are the averages for three sets of experiments; bars are one standard deviation. The pKa values are marked on the pH axis.

of exposed and matched chain lengths.30 In that study, Ru(NH3)5(4-AMP) redox centers were connected to a HS(CH2)15COOH SAM. In the matched case, the diluent thiols were HS(CH2)15COOH and the mean rate constant was 1 s-1. In the exposed case, the diluent thiols were HS(CH2)11COOH and the mean rate constant was 20 s-1. The change in standard rate constants was attributed to significant electronic coupling between the redox center and the electrode via the diluent thiols. Because the terminal group of the diluent thiol is different in the present data set, the factor of 10 change for the Os complex on the thicker and thinner SAM may be the same as the factor of 20 change for the Ru complex. Figure 9 shows the pH dependence of the reorganization energies of the OsIII and OsII forms in H2O, and Figure 10 contains the corresponding R(0) plot. Over nearly the entire pH range, the OsIII reorganization energy (0.6-0.7 eV) is roughly constant. Surprisingly, the OsII reorganization energy (0.9 ( 0.1 eV) is significantly larger than the OsIII reorganization energy. Figure 10 does not exhibit the predicted oscillation of R(0) for the stepwise model; see the Supporting Information. Values for R(0) hover around 0.46 ( 0.02 over most of the pH range, including pH values below pKa1. Only at pHs above pKa2 do the reorganization energies appear to converge to a common value of 0.7 eV and R(0) appear to shift toward 0.5. Similar behavior was observed on the thinner SAMs.29 Closer inspection of Figure 9 shows that the OsII and OsIII λ values approach each other at pHs near both of the pKa values and that the greatest difference between the two reorganization energies is at the intermediate pHs. This pattern can partly explain the pH dependence of the standard rate constant. Equation 1

Langmuir, Vol. 22, No. 25, 2006 10649

Figure 11. log(ks) vs pH or pD for the thicker SAM in H2O (squares) and D2O (triangles). Average values of three data sets (H2O) or four data sets (D2O) and associated standard deviation bars are shown. The solid lines are a guide for the eye. The two sets of pKa values are marked on the pH or pD axis.

Figure 12. Reorganization energies vs pD in D2O. Open squares, λ for OsIII; open diamonds, λ for OsII. The open symbols are the average values for four sets of experiments; bars are one standard deviation. For comparison, the solid circles are λ for OsIII in H2O, and the solid triangles are λ for OsII in H2O. The solid and dashed lines are a guide for the eye. The two sets of pKa values are marked on the pH or pD axis.

shows that higher reorganization energies for both species correspond to a lower standard rate constant. Similarly, the higher reorganization energies of OsII correspond to a decrease in the standard rate constant in the pH range of 4-8. Attempts to predict the changes in the standard rate constant using the analysis in the Supporting Information show that the measured reorganization energies do not quantitatively account for all of the pH dependence of log(ks). Data scatter also prevents a more quantitative comparison. However, it appears that part or all of the pH dependence of log(ks) arises from the pH dependence of λ for OsII. E. Kinetic Parameters in D2O. In Figure 11, the log(ks) vs pD plot in D2O is compared with the log(ks) vs pH plot in H2O on the thicker SAM. Above pH 4, the standard rate constants in D2O are approximately a factor of 2 slower than the standard rate constants in H2O. Below pH 4, the standard rate constants are the same within experimental uncertainty. The surprisingly distinct break in the log(ks) plot on D2O between pD 3 and 4 is under continuing investigation, but appears to be real. Overall, the kinetic isotope effect is small. The reorganization energies in D2O exhibit the same pattern with respect to pD as the reorganization energies in H2O (Figure 12). A plot of R(0) vs pD is very similar to the plot in Figure 10. The reorganization energies for the OsIII species in D2O are also very similar to the values obtained in H2O, but the reorganization energies for the OsII species are noticeably larger

10650 Langmuir, Vol. 22, No. 25, 2006

Madhiri and Finklea

Table 2. Comparison of Z Factors for Several Redox Molecules redox molecule Ru(NH3)5(py) Os(bpy)2(py)(Cl) Os(bpy)2(py)(H2O) Os(bpy)2(py)(H2O) Os(bpy)2(py)(D2O)

pH 5 3 11 11

ks (s-1)

average λ (eV)

FCb

0.9 0.6 0.75 0.65 0.8

2.2 × 10 4.9 × 10-4 1.0 × 10-4 2.9 × 10-4 6.2 × 10-5

1 10 1.5 4 1.5

-5

Z/104 (s-1)

ref

4.5 2.0 1.5 1.4 2.4

30 29 this work this work this work

a These complexes were attached to gold via HS(CH2)15C(O)N(H)CH2- in a SAM containing HS(CH2)15X (X ) CH2OH or COOH) as the diluent. Average reorganization energies have an uncertainty of (0.1 eV. b FC ) Franck-Condon factor is the integral and normalization factor in eq 1.

around pH 4-8. Also, OsII and OsIII λ values appear to converge at pH above pKa2 to a higher value in D2O than in H2O. A comparison of Figures 11 and 12 shows that the decrease of the standard rate constant in D2O is correlated with the larger reorganization energy for OsII in D2O relative to the value in H2O. F. Discussion. We reiterate our conclusion that the stepwise model is eliminated from consideration. This model predicts standard rate constants and R(0) values that are independent of pH at pH less than pKa1 and pH greater than pKa2.27 For the osmium complex, R(0) should be close to 0.5 in those pH domains. The model also predicts that R(0) oscillates below and then above 0.5 in the intermediate pH range with increasing pH. The data do not exhibit either behavior (see the Supporting Information). Even more convincing is the absence of a deep minimum in the plot of log(ks) vs pH. Assuming a reorganization energy of 0.9 eV for all four osmium species and a standard rate constant of 4.5 s-1 for the two half reactions (Figure 2), the stepwise model predicts that the minimum rate constant should be 60 times smaller than the rate constant at the pH extremes. The results in H2O show a decrease of only a factor of 4 (see the Supporting Information). Can concerted PCET account for the observed behavior of the osmium system? Concerted PCET can be described as the diagonal of the square scheme (Figure 2). The rate-determining step (RDS) is

OsIII(OH) + e + H+ T OsII(OH2)

(6)

However, this equation is misleading. Rate expressions based in eq 6 predict a cathodic rate constant at constant potential that is first-order on proton concentration and an anodic rate constant that is independent of pH. The data do not obey either prediction. The reaction partner for proton transfer is likely to be water at all pHs because of its high concentration relative to H+ and OH-. Consequently, the reduction step is better represented by

OsIII(OH)‚‚‚HOH + e f OsII(OH2)‚‚‚OH-

(7)

In the reactant, a water molecule forms a strongly hydrogenbonded partner with the hydroxide ligand. Because there is a finite delay between the reduction and oxidation steps even in the fast scan CVs, it is likely that the hydroxide product in eq 7 is replaced with water prior to the scan in the reverse direction. Consequently, the rate-determining anodic step is

OsII(OH2)‚‚‚OH2 f OsIII(OH)‚‚‚HOH2+ + e-

(8)

Again, water is the proton-transfer partner in the reactant. A major consequence of these rate-determining steps is that the pH does not appear explicitly in the rate law, so the rate of each reaction is, in the first approximation, independent of pH. However, the standard rate constants for the Os complex does show a pH dependence (Figures 8 and 11), with a minimum

value roughly midway between the two pKa’s, but with no change in slope at the pKa values in the plots. Another source of pH dependence might be the presence of alternative rate-determining steps at low (pH < pKa1) and high pH (pH > pKa2). At low pH, the rate-determining step should involve the fully protonated species with no proton transfer

OsIII(OH2)‚‚‚OH2 + e- T OsII(OH2)‚‚‚OH2

(9)

An equivalent expression can be written at high pH in which the ligand is OH. If this step has a different standard rate constant, then the data in Figures 8 and 11 might be based on two competing paths at low and high pH. However, there is a problem with the latter suggestion. The half-reaction in eq 9 is equivalent to the bottom step in the square scheme in Figure 2. For both paths, the standard rate constant should be constant and R(0) should be near 0.5. The data do not fit these expectations. The inverse correlation between the standard constants vs pH or pD (Figure 11) and the changes in λ for OsII (Figure 12) implies that this one parameter may be the key to the kinetic behavior of this 1 electron, 1 proton system. We do not have an explanation for the pH dependence of λ for OsII in H2O nor its more pronounced pH dependence in D2O. Even more puzzling is the fact that λ for OsII is larger than λ for OsIII. From electrostatic arguments, one would expect the bond force constants to be higher for OsIII(OH) than for OsII(OH2). Hence, the inner sphere reorganization energy for the OsIII species should be larger. The outer sphere reorganization energy, including both electron- and proton-transfer components, is anticipated to be nearly the same for both oxidation states, so the total reorganization energy of OsIII should be larger than the total reorganization energy for OsII. This argument was made by Hupp and Weaver to explain the kinetic behavior (asymmetrical Tafel plots, R(0) greater than 0.5) of aquated Cr2+/3+ and Eu2+/3+ ions.31 A theoretical explanation of these observations is highly desirable. The observation that the different λ values persist at pH less than pKa1 and greater than pKa2 (domains in which both mechanisms predict simple electron transfer) indicates that this behavior is not necessarily connected to proton-coupled electron transfer. When the two reorganization energies are nearly the same (at pH 3 and 11), it is instructive to compare the Z factor from eq 1 for various redox couples. Table 2 contains results for five redox couples which share a common tether to the gold electrode and two similar diluent thiols in the SAM. Within the uncertainties of the calculations, the Z factors for the osmium complex exhibiting the proton-coupled electron transfer are identical to the Z factors for redox molecules exhibiting simple electron transfer. Clearly, the electronic coupling between the electrode and the redox molecule is the dominant factor in determining Z. Quantitative calculation of the kinetic isotope effect via the ratio of Z is complicated by the differing reorganization energies for the OsII and OsIII forms. However, it appears that the kinetic isotope effect is less than a factor of 2.

Proton-Coupled Electron Transfer at a SAM

Summary The model 1e1H Os system attached to the longer SAM provides an opportunity to thoroughly explore potential, pH, and isotope effects on rate constants for both oxidation and reduction. The results confirm the earlier observation that the kinetics contradict the clear predictions of the stepwise mechanism of proton-coupled electron transfer. The evidence for the concerted mechanism is mainly the kinetic isotope effect. In any event, three observations need an explanation: (a) the weak but measurable pH dependence of the standard rate constant; (b) the persistent asymmetry of the Tafel plots and the resulting anomalous disparity of reorganization energies for the oxidized and reduced forms of the complex over the entire pH range; and (c) the change in the standard rate constant in D2O between pD 3 and 4. The changes in standard rate constants are inversely correlated with the pH-dependent reorganization energy of the OsII species. The Z factor is largely determined by electronic coupling between the metal electrode and the redox molecule and is approximately independent of pH. The kinetic isotope effect (ratio of Z factors in H2O and D2O) appears to be less than 2 at all pHs. In this case, with relatively weak coupling between the site of electron transfer (the osmium) and the proton-transfer site (the water ligand), the results oppose the stepwise PCET mechanism. A different picture is observed in the galvinol/ galvinoxyl system.26,28 This phenol-like redox molecule exhibits a wide separation of the two pKa values, and only pKa2 is experimentally accessible. The trends in log(ks) vs pH and R(0) vs pH do match the predictions of the stepwise model for pH greater than 7. However, at lower pH values, that system exhibits a deviation in the log(ks) values above the predicted trend. The deviation was very tentatively attributed to the onset of a concerted path.

Langmuir, Vol. 22, No. 25, 2006 10651

It is apparent that PCET systems are still in need of extensive experimental exploration to resolve the numerous anomalies and to guide theoretical treatments. These types of experiments should be extended to more biologically relevant redox systems, such as the quinones42-46 and catechols. A potential benefit may be the rational design of electrocatalysts for enhancing the rate constants for PCET. Such catalysts may find uses in the analyses of selected neurotransmitters (e.g., dopamine) during in situ monitoring in living brains.47 Acknowledgment. We thank Dr. Novruz Akhmedov for assistance with the NMR spectra. Supporting Information Available: Data interpretation for different reorganization energies of the oxidized and reduced forms; synthesis and characterization of [OsII(bpy)2(4-AMP)(H2O)](PF6)2; analysis of linear scan voltammograms (including simulations), comparison of the kinetic data with simulations based on the stepwise model, and thermodynamic parameters from the pH dependence of the formal potential. This material is available free of charge via the Internet at http://pubs.acs.org. LA061103J

(42) Mukae, F.; Takemura, H.; Takehara, K. Bull. Chem. Soc. Jpn. 1996, 69, 2461-2464. (43) Mo, Y.; Sandifer, M.; Sukenik, C.; Barriga, R. J.; Soriaga, M.; Scherson, D. Langmuir 1995, 11, 4626-4628. (44) Hong, H.-G.; Park, W. Langmuir 2001, 17, 2485-2492. (45) Kazemekaite, M.; Bulovas, A.; Talaikyte, Z.; Butkus, E.; Railaite, V.; Niaura, G.; Palaima, A.; Razumas, V. Tetrahedron Lett. 2004, 45, 3551-3555. (46) Larsen, A. G.; Gothelf, K. V. Langmuir 2005, 21, 1015-1021. (47) Bath, B. D.; Martin, H. B.; Wightman, R. M.; Anderson, M. R. Langmuir 2001, 17, 7032-9.