J. Phys. Chem. B 2001, 105, 8739-8745
8739
Redox Kinetics in Monolayers on Electrodes: Electron Transfer Is Sluggish for Ferrocene Groups Buried within the Monolayer Interior† James J. Sumner and Stephen E. Creager* Department of Chemistry, Clemson UniVersity, Clemson, South Carolina 29634-0973 ReceiVed: April 2, 2001; In Final Form: June 16, 2001
A study was undertaken of the redox chemistry of ferrocene groups in alkanethiolate-based monolayer films on gold electrodes. The monolayers are structured so as to position the ferrocene groups within the interior portion of the monolayer. The study focused on the difference in the redox thermodynamics (oxidationreduction potential) and kinetics (standard electron-transfer rate constant) for ferrocene groups at the monolayer-solution interface and those in the monolayer interior. Ferrocene groups that were buried in the monolayer interior exhibited oxidation-reduction potentials that were strongly shifted to positive potentials relative to the potentials found for similar monolayers in which the ferrocene groups are exposed to electrolyte solution, in agreement with prior work on similar systems. The principal new finding from this study is that the redox kinetics associated with ferrocene oxidation/reduction were much more sluggish when the ferrocene groups were buried within the monolayer interior when compared with monolayers in which the ferrocene groups were exposed to electrolyte. For example, the standard electron-transfer rate constant for ferrocene oxidation/reduction in a monolayer of ferrocenyl-decanethiol coadsorbed with decanethiol was found to be approximately 40 000 s-1, whereas that for a monolayer of ferrocenyl-decanethiol coadsorbed with 1-mercaptoeicoscane (C20H41SH) was approximately 200 s-1. The differences in behavior for “buried” and “exposed” ferrocene groups were found to depend critically upon the electrolyte anion size, i.e., ferrocene oxidation/ reduction was almost completely inhibited when sodium poly(p-styrene sulfonate) was used as the electrolyte. These results are interpreted in terms of a reaction involving coupled electron- and ion-transfer, i.e., electrontransfer from ferrocene to the electrode much occur concomitantly with anion transfer from the electrolyte to the oxidized ferricenium site in the monolayer, and the rate of electron transfer can depend critically upon the rate of ion transfer into the monolayer.
Introduction The role of counterion motion in transporting electronic charge in redox-active assemblies is widely recognized.1 For example, it has long been known that the nature and concentration of the electrolyte can have strong effects on the voltammetric wave shape for oxidation/reduction of many types of redox-active films on electrodes, including redox polymers, electronically conductive polymers, and redox-active inorganic materials such as Prussian Blue. A particularly striking example of such an effect was noted by Elliott and co-workers when they showed that certain redox reactions in redox polymer films on electrodes could be suppressed by choosing an electrolyte for which one type of ion (e.g., the cation for a film undergoing reduction) is so large that it cannot penetrate the film interior.2,3 Other workers have noted the critical role of counterion motion in establishing concentration gradients of oxidized and reduced sites in redox-active assemblies. Mixed-valent materials can be made by partial oxidation/reduction of redox-active films, and the resulting concentration gradients can be frozen in place by changing the conditions such that counterion motion is suppressed.4,5 Such films can retain their frozen-in concentration gradients (or lack thereof) even when a high electric field is applied that can drive electron transport through the film by field-driven electron hopping between sites. These works and many related works serve to demonstrate the crucial role of †
Part of the special issue “Royce W. Murray Festschrift”. * Corresponding author.
coupling between electron and ion transport in materials that can transport both kinds of charge carriers.1,6,7 The role of ion motion in the transport of charge within redoxactive monolayer assemblies has also been considered, albeit in less depth. Creager and Rowe8-11 and also Uosaki and coworkers,12,13 have noted that the apparent formal potential for ferrocene oxidation in alkanethiolate-based self-assembled monolayers on gold electrodes depends strongly on the nature and concentration of the electrolyte, particularly the anion. A crucial role for ion-pairing between the electrogenerated ferricenium and the electrolyte anion was proposed as the cause of this effect. Other workers have noted similar effects in related monolayer assemblies.14-21 Smith and White22 and later Fawcett and co-workers23,24 and Honeychurch and co-workers25 have noted that a net separation between charged sites in a monolayer and charge-compensating counterions in solution can bring about an apparent shift in the formal potential for the surface-confined redox reaction. Mirkin and co-workers reported that the redox reactions of azobenzene derivatives in monolayers are strongly affected depending upon whether the azobenzene units were exposed to the solution or buried in the monolayer.26 Azobenzene units that were buried in the monolayer were found to be immune from reduction even when the applied potential was sufficient to bring about oxidation or reduction. These authors termed this effect “ion-gated electron transport”, in recognition of the fact that oxidation/reduction is strongly affected by the nature of the medium surrounding the redox-active agents, and
10.1021/jp011229j CCC: $20.00 © 2001 American Chemical Society Published on Web 08/10/2001
8740 J. Phys. Chem. B, Vol. 105, No. 37, 2001 specifically upon the access of ions to the redox-active site. Of course, in the case of azobenzene, oxidation/reduction involves both electron and proton transfer, and any effects on reactivity must be considered in light of the local pH in the monolayer as well as any intrinsic effects of the monolayer local microenvironment on the redox reactivity. In the present work, we consider the redox kinetics for ferrocene groups in two different local microenvironments at monolayer-coated electrodes. In one case, the ferrocene groups are exposed to the electrolyte solution at the monolayer-solution interface by making the monolayer out of ferrocene-capped and nonferrocene-capped alkanethiol chains of equal length. In another case, the ferrocene groups are “buried” within the monolayer interior by making the monolayer out of the same ferrocene-capped alkanethiols used to make the exposed monolayers, but with nonferrocene-capped alkanethiol coadsorbates that have much longer alkane chains. Figure 1 illustrates the four monolayer structures in question. (We note that the structures in Figure 1 are for illustration purposes only and are not meant to indicate that the alkane chains actually exist in the conformations shown. In fact, it seems likely that the alkane chains surrounding the buried ferrocene groups are probably disordered, although direct evidence for this is difficult to obtain since the monolayers in question have a very low surface coverage of ferrocene, and the response in most conventional surface analytical experiments (e.g., reflection infrared spectroscopy) would probably be dominated by the portions of the monolayer that are not involved in burying the ferrocene groups.) These studies are important because they can provide insight into the effect of the local microenvironment surrounding a redox-active molecule on the dynamics of charge transport between the redox molecule and an electrode. This in turn is important in understanding the factors that must be considered when designing molecular-scale electronic devices that transport charge within organized molecular assemblies in such a way as to accomplish specific functions, and also in understanding the behavior of redox-active cofactors in biological structures such as redox proteins. The key finding from this work is that the apparent electrontransfer rate between ferrocene and an underlying gold electrode is much slower for the case of buried ferrocene groups than for the case of exposed ferrocene groups. A proposed explanation for this effect is that it is caused by slow ion motions between the bulk electrolyte and the interior of the monolayer for the case of the buried ferrocenes.
Sumner and Creager
Experimental Section
Figure 1. Illustrative structures of ferrocene-containing alkanethiolate monolayers: (A) Fc-C12H24-SH/C12H25-SH; (B) Fc-C10H20-SH/ C10H21-SH; (C) Fc-C12H24-SH/C22H45-SH; and (D) Fc-C10H24SH/C20H41-SH.
Materials. Gold wire for electrodes was obtained from Alfa Aesar (0.127 mm diameter, Premion grade, >99.99% pure). Ethanol (100%) was from AAPER Alcohol and Chemical Company. Perchloric acid (70% in water, reagent ACS grade) and p-toluene sulfonic acid in the sodium salt form were obtained from Acros. Water was purified with a NANOpure water filtration system (Barnstead) to a resistivity of at least 12 Mohm cm. Poly(sodium 4-styrenesulfonate) (M.W. ∼ 70 000), 1-mercaptodecane, 1-mercaptododecane (98+%), 1-bromodocosane (C22H45-Br) and 1-bromoeicoscane (C20H41-Br) were all obtained from Aldrich. All reagents from commercial sources were used as received. The ferrocenyl-alkanethiols and the alkanethiols used to prepare monolayers were synthesized as described previously.27 Electrode and Monolayer Preparation. The gold electrodes were prepared by first melting the tip of the gold wire in a
Bunsen burner flame, forming a small sphere. The diameter of the sphere was approximately 0.50 mm. Mixed monolayers were formed on the electrode surface by soaking in a coating solution of thiols in ethanol. The coating solutions were typically 1 mM total thiol concentration at a 1:10 ratio of ferrocenyl-thiol to coadsorbate thiol, though the ratio was sometimes adjusted as needed to achieve the desired ferrocene surface coverage in the monolayers. The electrodes were exposed to this coating solution for 10 minutes, rinsed with water, then rinsed with 2-propanol and placed in a 500 mM ethanolic solution of the coadsorbate thiol at ∼ 45 °C for at least thirty minutes. The electrodes were then cooled to room temperature on the benchtop slowly (over two or more hours). The heat cycling step in the coadsorbate solution was performed to displace any weakly bound ferrocenyl-alkanethiols and fill in any pinholes in the monolayer to provide a more stable and structured system.
Redox Kinetics in Monolayers on Electrodes
J. Phys. Chem. B, Vol. 105, No. 37, 2001 8741
Figure 2. Comparison of ac voltammograms of mixed monolayers of (A) Fc-C12H24-SH/C12H25-SH at 5, 50, and 500 Hz as well as (C) FcC12H24-SH/C22H45-SH at 5, 50, and 500 Hz. Note that the peak size remains relatively constant in the exposed ferrocene system (A) through 50 Hz but dramatically decreases in the buried ferrocene system (C) after 5 Hz.
Electrochemical Characterization. After rinsing with 2-propanol and then with water, the monolayer-coated electrode was placed in a three-electrode cell containing an aqueous electrolyte, a silver/silver chloride/saturated potassium chloride reference electrode and a platinum wire counter electrode. The working electrode was positioned so that the gold sphere was just below the surface of the electrolyte, with as little of the wire exposed to the solution as possible. The cell time constant in this configuration is estimated as follows. The electrode area estimated from the radius of the gold sphere is approximately 2 × 10-3 cm2, the specific capacitance of the alkanethiol-coated gold is approximately 2 µF cm-2, and the solution conductivity for a 1.0 M HClO4 solution is approxamately 340 mS cm-1, thus the interfacial capacitance is approximately 4 nF and the solution resistance is approximately 10 Ω which leads to an estimated cell time constant of approximately 40 ns. Electrical connection was made at the end of the wire opposite the sphere. Redox kinetics for ferrocene groups on the coated electrodes were characterized using a recently described variant of the ac voltammetry method.28 The ac voltammograms were collected on a Solartron SI 1260 impedance/gain-phase analyzer connected to a Solartron SI 1287 electrochemical interface. Both were connected to a PC by GPIB for control and data collection using Zplot and Zview software by Scribner Associates, Inc. Results AC voltammetry has proved to be particularly useful for investigating the dynamics of redox reactions in monolayer assemblies on electrodes.28-33 For a reversible reaction, the position of the peak in the voltammogram provides information about the formal potential of the surface redox reaction, and the magnitude of the ac peak current provides information about the amount of redox-active material on the electrode.34 Furthermore, changes in the magnitude of the ac peak current with the changing frequency of the applied ac voltage provide information about the rate of electron transfer between the redox molecules and the electrode.28
Figures 2 and 3 present a series of ac voltammograms at selected frequencies which illustrate the behavior of the “exposed” ferrocenes (monolayers A and B) and “buried” ferrocenes (monolayers C and D) in the present monolayer series. Two aspects of these voltammograms merit comment. First, the peak potentials for monolayers with buried ferrocenes are always shifted in a positive direction relative to those for monolayers with exposed ferrocenes. For example, in Figure 2 the peak potential for the ferrocene group in monolayer A is approximately +0.37 V vs Ag/AgCl at 50 Hz, whereas the peak potential for ferrocene in monolayer C is approximately +0.60 V at the same frequency. This corresponds to a positive shift of 230 mV. A similar shift was noted for monolayer D when compared with monolayer B (Figure 3). These shift are reminiscent of similar shifts that were noted in earlier work with similar monolayers studied using slow-scan cyclic voltammetry.8-11 In the earlier work the shifts had been attributed to a combination of local solvation effects on the ferrocene/ ferricenium redox potential, and double-layer effects whereby the apparent redox potential is shifted positive due to a net charge separation between the oxidized redox molecules and their charge-compensating counterions in solution. It seems likely that a similar explanation applies in the present case. The second important aspect of these voltammograms that merits comment is the gradual diminution of the peak currents relative to the background as frequency increases. This effect has been attributed in earlier work to the fact that as the frequency increases, the redox reaction eventually cannot keep pace with the rapidly fluctuating potential and thus contributes less to the periodic ac current.28-31 In the present case, the critical aspect of this effect is that it occurs over a different frequency regime for the “exposed” ferrocenes than it does for the “buried” ferrocenes. This difference can be seen particularly clearly in Figure 3 for monolayers B and D: for monolayer B with the exposed ferrocene group, an increase in frequency from 100 to 1000 Hz has almost no effect on the peak height relative to baseline, whereas the same change in frequency for mono-
8742 J. Phys. Chem. B, Vol. 105, No. 37, 2001
Sumner and Creager
Figure 3. Comparison of ac voltammograms of the mixed monolayers of (B) Fc-C10H20-SH/C10H21-SH at 100 and 1000 Hz as well as (D) Fc-C10H20-SH/C20H41-SH at 10, 100, and 1000 Hz. Note the peak size remains relatively constant in the exposed ferrocene system (B) through 1000 Hz, but decreases in the buried system (D) after 10 Hz.
layer D with the buried ferrocene causes the peak to almost disappear. In qualitative terms, this observation suggests that the redox kinetics are more sluggish for the buried ferrocenes in monolayer D than for the exposed ferrocenes in monolayer B. Figures 4 and 5 present a more quantitative assessment of the relative electron-transfer rates for the exposed and buried ferrocenes. The figures present plots of the ratio of the peak current to the background current (hereafter called the peak current ratio) against log(frequency) for a series of ac voltammograms spanning a wide frequency range. The plots all show the expected sigmoidal shape, with the plateau value of the peak current ratio at low-frequency being indicative of the amount of ferrocene on the electrode and the breakpoint frequency being indicative of the standard electron-transfer rate constant for the particular ferrocene derivative being studied. (Note that the cell time constant of 40 ns estimated as described in the Experimental Section is much faster than any of the electron-transfer reactions characterized in this work, which enables us to rule out the cell time constant as a factor causing the observed diminution of ACV peak current with increasing frequency.) The figure legends and captions list specific values of the standard rate constants that were obtained from fits to the data (solid lines) using a Randles equivalent circuit model to describe the behavior at the peak and a simple series RC circuit to describe the behavior away from the peak. The trends in the fitted rate constants with respect to monolayer structure are quite clear; both monolayers with an exposed ferrocene group exhibit rate constants that are large and strongly dependent on the length of the alkane chain linking ferrocene to the electrode (i.e., ko ) 1700 s-1 for monolayer A and 40 000 s-1 for monolayer B), whereas both monolayers with a buried ferrocene exhibited rate constants that were much lower and approximately independent of the length of the chain linking ferrocene to the electrode (i.e., ko ) 200 s-1 for monolayers C and D). It is significant that the difference arises even for monolayer pairs (A and C, B and D) for which the nature and length of the molecular bridge linking the ferrocene group to the electrode is identical. To our
Figure 4. Ipeak/Ibackground versus log(frequency) plots of mixed monolayers of Fc-C12H24-SH/C12H25-SH (A, top) and Fc-C12H24-SH/ C22H45-SH (C, bottom). The lines represent fits to the data using a Randles equivalent circuit, as described in ref 28. Fitting parameters are as follows: (A), standard electron-transfer rate k0 ∼ 1700 s-1 with a surface coverage of 5.25 × 10-13 mol cm-2; (B) k0 ∼ 200 s-1 with a surface coverage of 3.65 × 10-13 mol cm-2.
Redox Kinetics in Monolayers on Electrodes
Figure 5. Ipeak/Ibackground versus log(frequency) plots of mixed monolayers of Fc-C10H24-SH/C10H21-SH (B, top) and Fc-C10H20-SH/ C20H41-SH (D, bottom). The lines again represent fits to the data using a Randles equivalent circuit. Fitting parameters are: (B) k0 ∼ 40 000 s-1 with surface coverage of 2.20 × 10-12 mol cm-2; and (D) k0 ∼ 200 s-1 with surface coverage of 9.3 × 10-13 mol cm-2.
knowledge, this is the first definitive demonstration of a large local microenvironmental effect on the electron-transfer rate (as distinguished from the redox potential) in a redox-active monolayer assembly on an electrode surface. The data in Figures 2-5 were not particularly difficult to obtain, however it was necessary to restrict the study to monolayers in which ferrocene groups were present only in very small amounts. If the ferrocene coverage was too large, then it was found that the frequency dependence of the acv peak currents could no longer be fit by a simple Randles equivalent circuit. This point is illustrated in Figure 6, which compares the peak-current-ratio plots for two monolayers, one with an exposed ferrocene (monolayer A) and another with a buried ferrocene (monolayer C) present at relatively high surface coverages. The trend in the data is clear; the behavior of the exposed ferrocene system is still well described by a simple Randles circuit with a single-valued rate constant even at a relatively high coverage, whereas the behavior of the buried system has become complex and can no longer be fit using a simple Randles circuit at high coverage. In fact, behavior very similar to that found for the buried-ferrocene system at high coverage has been reported in simulations using a modified Randles circuit that accounts for the presence of multiple populations of redox sites with different rate constants.35 There is not enough information at hand in the present data set to enable a unique fit to a particular rate constant distribution; however, even in the absence of such a fit, the fact that the
J. Phys. Chem. B, Vol. 105, No. 37, 2001 8743
Figure 6. Ipeak/Ibackground versus log(frequency) plots for mixed monolayers of Fc-C12H24-SH/C12H25-SH (A, top) and Fc-C12H24-SH/ C22H45-SH (C, bottom) at relatively high surface coverages of the redox species. The lines again represent fits to the data using a Randles equivalent circuit. Fitting parameters are as follows: (A) k0 ∼ 1800 s-1 with a surface coverage of 7.60 × 10-12 mol cm-2; (C) k0 ∼ 200 s-1 with a surface coverage of 1.30 × 10-11 mol cm-2. Note that the quality of the fit for the simluation is good for the “exposed” system (A), but not for the “buried” system (C).
behavior of the buried system at high coverage is qualitatively similar to that expected from a system with multiple rate constants suggests that, at high coverage, the monolayer structure is such that multiple microenvironments exist for ferrocene groups, and that those different microenvironments are manifest as a distribution in electron-transfer rate constants. The data in Figures 2-5 all indicate that the apparent electron transfer rates for buried ferrocene groups are suppressed relative to the rates for the exposed ferrocenes, but they offer no insight into why this might be so. One possibility is that the rate is coupled to the rate of transport of charge-compensating counterions in to and out of the monolayer. If ready access to counterions is required for electron transfer to occur, and if counterion transport is sluggish, then the overall rate could be set not by the rate of the electron transfer event itself, but rather by the rate of counterion transport. One way to examine this hypothesis is to study the effect of changing the electrolyte anion size on the apparent electron-transfer rates for exposed and buried ferrocenes. Figure 7 presents an ac voltammetry data set for monolayers A (exposed) and C (buried) that was acquired to conduct such a test. In these experiments the usual 1.0 M perchloric acid electrolyte was replaced first with a 0.1 M solution of sodium poly(p-styrene sulfonate), and then with a mixed solution containing both 0.1 M sodium poly(p-styrene
8744 J. Phys. Chem. B, Vol. 105, No. 37, 2001
Sumner and Creager present, but it increases in magnitude by more than 10 times when the monoanionic electrolyte is added. We attribute this effect to the p-toluenesulfonate anions being able to penetrate the monolayer interior much more effectively than poly(styrene sulfonate), thereby “turning on” the ferrocene oxidation reaction that was otherwise inhibited due to a lack of access of anions to the redox-active sites. The polymeric poly(styrene-sulfonate) anions are too large to penetrate the monolayer interior, and consequently the ferrocene oxidation is inhibited when they are the only anions present in solution. Discussion
Figure 7. Comparison of 10 Hz ac voltammograms of the mixed monolayers of Fc-C12H24-SH with exposed redox sites (top, monolayer A) and buried redox sites (bottom, monolayer C) with polymeric electrolyte (0.1 M sodium poly(4-styrene sulfonate)) and polymeric electrolyte doped with monomeric electrolyte (0.1 M sodium poly(4styrene sulfonate) and 0.1 M sodium p-toluene sulfonate).
sulfonate) and 0.1 M sodium p-toluene sulfonate. This pairing of electroltye anions was chosen because the chemical nature of the anions is similar in both cases, but in the case of the polyanionic electrolyte the anions are very large and sterically bulky, such that anion transport into the interior of the alkanethiolate monolayer should be inhibited. Thus, a comparison of the behavior in the presence of only the polyanionic electrolyte with that in the presence of polymeric and monomeric electrolyte reveals something about the effect of anion size on the redox behavior. In all cases, the use of p-toluene sulfonate and/or poly(pstyrene sulfonate) in place of perchlorate in the electrolyte produces a voltammetric peak potential that is shifted in a positive direction relative to that obtained in 1.0 M HClO4. This effect is probably attributable to the different degrees of ion pairing for perchlorate and the aryl sulfonate anions with ferricenium, as was noted in earlier work using cyclic voltammetry to study redox potentials in shorter-chain ferrocenecontaining monolayers with similar structures.8-11 For the case of the exposed ferrocene (monolayer A) the difference in behavior between the polyanionic electrolyte without and with the monoanionic electrolyte present is relatively minor, consisting only of a small increase in the magnitude of the peak current when the monoanionic electrolyte is added. In contrast, the effect is much larger for the buried ferrocene (monolayer C); in that case, the ferrocene voltammetric peak is nearly indistinguishable from background when only the polyanionic electrolyte is
The effects of ion-pairing on the dynamics of electron transport over long distances have previously been considered. Marcus has recently presented an excellent summary of some recent literature on this topic.36 In his paper, Marcus considers three possible scenarios by which ion-pairing interactions could affect redox kinetics in bridged donor-acceptor molecules. Those three scenarios are (i) electron transfer followed by ion transfer; (ii) ion-pair dissociation followed by electron transfer; and (iii) ion transfer accompanying electron transfer. The first two scenarios could be considered as special cases of redox reactions with either a following or a preceding chemical step. In the latter scenario, the ion association/dissociation steps are thought to contribute to the reorganization energy associated with electron transfer. In all cases, Marcus showed that the apparent electron-transfer rate constants may be considered in terms of a reciprocal sum of rate constants for the electrontransfer event and the ion-transport event, the latter often consisting of a combination of an association or dissociation rate constant and an ion-pairing equilibrium constant. The details of how these rate constants and equilibrium constants combine are different for each specific scenario, but the overall trend is similar in all cases; a slow ion-transport step in a system involving coupled electron and ion transport can act to limit the overall charge-transport rate. To our knowledge, no similar treatment has been made of the effect of ion-pairing interactions in long-range electrontransfer reactions in monolayer assemblies on electrodes. There are some critical differences between such reactions and the intermolecular and intramolecular electron-transfer reactions that Marcus considers in his work. One critical difference is that the monolayer-based electrode reaction is really a half-cell reaction, and ion-pairing interactions are likely to be important only for the redox molecules (and then only for the charged forms) and not for the electrode. This is in contrast with the case of molecular electron donor-acceptor systems (bridged or otherwise), where each participant in the reaction must always be charged in at least one redox state. There may be other critical differences as well, and a quantitative treatment of the diminution of apparent electron-transfer rate constants for the buried ferrocene compared with the exposed ferrocene systems found in the present work is probably premature in the absence of at least a putative theoretical framework. Even so, the present observation that apparent electron-transfer rates are relatively sluggish in the buried-ferrocene systems, particularly when the anions are very large and sterically bulky, combined with the insights offered by Marcus that diminished electron-transfer rates can be caused by slow ion transport to and from charged sites involved in long-range redox reactions, does provide at least a qualitative understanding of the effect. Summary A study of the redox chemistry of ferrocene groups buried within the interior of alkanethiolate-based self-assembled mono-
Redox Kinetics in Monolayers on Electrodes layer films on gold electrodes was undertaken. The study showed that the rate of ferrocene oxidation/reduction was dramatically slowed for ferrocene groups in the monolayer interior compared with the rates for ferrocene groups exposed to the electrolyte solution. This was true even though the nature and length of the molecular bridge linking ferrocene to gold is the same in both cases. The finding is interpreted in terms of a reaction involving coupled ion and electron transfer to accomplish ferrocene oxidation/ferricenium reduction in the monolayers. References and Notes (1) Molecular Design of Electrode Surfaces; Techniques of Chemistry Series; Murray, R. W., Ed.; John Wiley & Sons: New York, 1992; Vol. 22. (2) Elliott, C. M.; Redepenning, J. G.; Balk, E. M. J.f Electroanal. Chem. 1986, 213, 203-215. (3) Elliott, C. M.; Redepenning, J. G.; Balk, E. M. J. Am. Chem. Soc. 1985, 107, 8302-8304. (4) Jernigan, J. C.; Chidsey, C. E. D.; Murray, R. W. J. Am. Chem. Soc. 1985, 107, 2824-2826. (5) Jernigan, J. C.; Murray, R. W. J. Phys. Chem. 1987, 91, 20302032. (6) Saveant, J. M. J.f Electroanal. Chem. 1988, 242, 1-21. (7) Saveant, J. M. J. Phys. Chem. 1988, 92, 1011-1013. (8) Rowe, G. K.; Creager, S. E. J. Phys. Chem. 1994, 98, 5500-5507. (9) Rowe, G. K.; Creager, S. E. Langmuir 1991, 7, 2307-2312. (10) Creager, S. E.; Rowe, G. K. J. Electroanal. Chem. 1997, 420, 291299. (11) Creager, S. E.; Rowe, G. K. Anal. Chim. Acta 1991, 246, 233239. (12) Uosaki, K.; Sato, Y.; Kita, H. Langmuir 1991, 7, 1510-1514. (13) Shimazu, K.; Yagi, I.; Sato, Y.; Uosaki, K. J. Electroanal. Chem. 1994, 372, 117-124. (14) Redepenning, J.; Tunison, H. M.; Finklea, H. O. Langmuir 1993, 9, 1404-1407.
J. Phys. Chem. B, Vol. 105, No. 37, 2001 8745 (15) Redepenning, J.; Flood, J. M. Langmuir 1996, 12, 508-512. (16) Pilloud, D. L.; Chen, X. X.; Dutton, P. L.; Moser, C. C. J. Phys. Chem. B 2000, 104, 2868-2877. (17) John, S. A.; Kitamura, F.; Nanbu, N.; Tokuda, K.; Ohsaka, T. Langmuir 1999, 15, 3816-3822. (18) Zhang, J. J.; Lever, A. B. P.; Pietro, W. J. J. Chem. Soc., Faraday Trans. 1997, 93, 3355-3362. (19) Ohtani, M.; Kuwabata, S.; Yoneyama, H. Anal. Chem. 1997, 69, 1045-1053. (20) Forster, R. J.; Faulkner, L. R. J. Am. Chem. Soc. 1994, 116, 54445452. (21) Acevedo, D.; Abruna, H. D. J. Phys. Chem. 1991, 95, 9590-9594. (22) Smith, C. P.; White, H. S. Anal.Chem. 1992, 64, 2398-2405. (23) Fawcett, W. R. J. Electroanal. Chem. 1994, 378, 117-124. (24) Andreu, R.; Calvente, J. J.; Fawcett, W. R.; Molero, M. J. Phys. Chem. B 1997, 101, 2884-2894. (25) Honeychurch, M. J.; Rechnitz, G. A. Electroanalysis 1998, 10, 285-293. (26) Campbell, D. J.; Herr, B. R.; Hulteen, J. C.; Van Duyne, R. P.; Mirkin, C. A. J. Am. Chem. Soc. 1996, 118, 10211-10219. (27) Creager, S. E.; Rowe, G. K. J. Electroanal. Chem. 1994, 370, 203. (28) Creager, S. E.; Wooster, T. T. Analy. Chem. 1998, 70, 4257-4263. (29) Sumner, J. J.; Weber, K. S.; Hockett, L. A.; Creager, S. E. J. Phys. Chem. B 2000, 104, 7449-7454. (30) Creager, S.; Yu, C. J.; Bamdad, C.; O’Connor, S.; MacLean, T.; Lam, E.; Chong, Y.; Olsen, G. T.; Luo, J.; Gozin, M.; Kayyem, J. F. J. Am. Chem. Soc. 1999, 121, 1059-1064. (31) Sumner, J. J.; Creager, S. E. J. Am. Chem. Soc. 2000, 122, 1191411920. (32) Brevnov, D. A.; Finklea, H. O. J. Electrochem. Soc. 2000, 147, 3461-3466. (33) Nahir, T. M.; Bowden, E. F. J. Electroanal. Chem. 1996, 410, 9. (34) O’Connor, S. D.; Olsen, G. T.; Creager, S. E. J. Electroanal. Chem. 1999, 466, 197-202. (35) Li, J.; Schuler, K.; Creager, S. E. J. Electrochem. Soc. 2000, 147, 4584-4588. (36) Marcus, R. A. J. Phys. Chem. B 1998, 102, 10071-10077.
