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Coupled Electron/Proton Transfer Studies of Benzoquinone-Modified Monolayers Wenbin Zhang, Scott M. Rosendahl, and Ian J. Burgess* Department of Chemistry, UniVersity of Saskatchewan, Saskatoon, Saskatchewan, S7N 5C9 Canada ReceiVed: October 5, 2009; ReVised Manuscript ReceiVed: December 29, 2009
The electrochemical reduction of benzoquinone involves a coupled transfer of protons and electrons. Theoretical treatments of proton-coupled electron transfer (PCET) indicate that both thermodynamic (apparent formal potential) and kinetic (apparent standard heterogeneous rate constants) quantities should be highly pH-dependent. Although there are several reports of quinone-derivatized self-assembled monolayers (SAMs) in the literature, no systematic studies on the pH dependence of the PCET kinetics have been reported. In this work, we outline a method to produce SAM-aminobenzoquinone monolayers starting from diluted monolayers of an amine-terminated SAM. The electroactive monolayer behaves as a nearly ideal Nernstian system (no interactions between redox centers). In neutral to weakly acidic electrolytes, the aminobenzoquinone can be reduced in a kinetically slow, two-electron, two-proton (2e2H) process, but in lower pH solutions, the observed reaction is consistent with an overall 2e3H transfer. Kinetic measurements have been made using techniques that do not rely on the assumption of a potential-independent apparent transfer coefficient. The apparent standard rate constants range from 0.1 to 0.01 s-1, which is roughly an order of magnitude faster than previous reports for comparable systems. Plots of log(ks,app) versus pH are distorted variants of the expected W plots predicted from the theory of 2e2H PCET. These deviations may be caused by surface pKa effects and/or additional reaction pathways arising from the third proton transfer. 1. Introduction The redox behavior of hydroquinone/quinone redox couples are particularly interesting examples of proton-coupled electron transfer (PCET).1–7 In aqueous solutions, with the pH lower than the pKa of the phenolic protons (∼108), the oxidation of dihydroquinone to its corresponding quinone involves the transfer of two electrons (2e-) and two protons (2H+). The exact mechanism of PCET for free quinones in both buffered9 and unbuffered10–14 aqueous solutions has been extensively debated for decades, and the largely accepted mechanism involves four, sequential steps (H+,e-,H+,e-at low pH and e-,H+,e-,H+ above pH 6)5,6,15salthough some recent studies have explored the possibility of concerted proton-electron transfer.16–18 These most recent theoretical advancements have been reviewed by Hammes-Schiffer and Soudackov.19 In the case of sequential PCET, the thermodynamics and kinetics of quinone/hydroquinone charge transfer are predicted to be pH-dependent, and in a series of papers, Laviron developed a theoretical treatment of PCET that predicts the pH dependence observed in experimental studies.1–6 Quinones (or their hydroquinone analogues) can also be localized on surfaces by chemical modification of the terminus of a thiol self-assembled monolayer (SAM), either before selfassembly or post facto. Over the past ten years or so, many electroactive monolayer systems based on thiol SAMs terminated with quinone derivatives have been studied. Such systems are very appealing because they allow precise control of the localized environment and the distances involved in the charge transfer between the electron acceptor and donor. For example, functionalized SAMs have been used to evaluate the influence of electron delocalization along the spacer separating the metal electrode and quinone centers.20–23 Furthermore, a quinone* To whom correspondence
[email protected].
should
be
addressed.
E-mail:
terminated monolayer provides a versatile surface upon which additional, clean, and quantitative chemistry can be performed. Examples include Diels-Alder reactions,24–26 molecular wire constructs,27–29 and peptide immobilization via oxime conjugates.30 There have now been several fundamental studies of quinone-SAM electrochemistry in aqueous solutions.20–23,28,29,31–42 Hong and Park described very slow kinetics (apparent standard rate constant, ks,app, ) 3.6 × 10-4 s-1) for a benzoquinone/ hydrobenzoquinone terminus separated from the thiol-Au surface by a 12-carbon methylene chain.35 Similarly, Ye et al. reported very slow PCET for 2-(-11-mercaptoundecyl)hydroquinone SAMs, as evidenced by large peak separations in cyclic voltammograms (CVs) recorded in acidic electrolytes.41 Trammel et al. also saw comparably small apparent standard rate constants for a slightly shorter-chained (eight CH2 units) quinone-SAM,20 and others have reported similar values for long-chained anthroquinone (AQ)37 and pyrroloquinoline quinone (PQQ) derivatives.39 In contrast, Abhayawardhana and Sutherland have very recently reported a ks,app ∼ 10 s-1 for an AQ monolayer with a 10-carbon alkyl spacer.31 This is roughly 3 orders of magnitude faster than the reported apparent rate constants mentioned above. However, none of these studies have provided an in-depth analysis of the pH dependence on ks,app, which has been described theoretically. The classical treatment of 2e2H PCET proposed by Laviron1–5 assumes a potential invariant value of 1/2 for all transfer coefficients, R, associated with each electron-transfer step. This treatment leads to a potential- and pH-dependent apparent transfer coefficient, Rapp. Finklea has recently proposed a modern theory of PCET for electroactive monolayers that includes potential-dependent transfer coefficients,7 in accordance with Marcus theory.43 This revised treatment predicts more or less the same qualitative dependence on the apparent rate constants with pH, but with systematically lower values of ks,app. Although Finklea and Haddox were able to apply their treatment to the
10.1021/jp909540c 2010 American Chemical Society Published on Web 01/22/2010
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one-electron, one-proton coupling of galvinol/galvinoxyl,44,45 to the best of our knowledge, there have been no reports testing the validity of the model to a 2e2H example. In this paper, we report the kinetic and thermodynamic pH dependence of aminobenzoquinone-monolayers. Using methods first reported by Lukkari et al.,28 and later refined by Nova´k and co-workers,32,33 we are able to build a monolayer system where the aminobenzoquinone surface concentration is sufficiently dilute enough to provide nearly ideal electrochemical behavior. We have used both voltammetry and chronocoulometry to extract values of the apparent rate constant and apparent formal potential as a function of pH. 2. Experimental Section Materials and Reagents. All chemicals were used without purification, except for the 1,4-benzoquinone (Alfa Aesar), which was purified by sublimation at reduced pressure and slightly elevated temperatures (∼40 °C). Electrolyte solutions were made from NaClO4 (g99.0%, Fluka) and buffered using ACS purity Na2HPO4 and NaH2PO4 · H2O. More acidic solutions were obtained by the addition of HClO4 to a pH 3 sodium phosphate buffer solution. All electrolytes were made using 18.2 MΩ · cm Millipore water. To ensure that the ionic strength was kept constant, electrolytes consisted of 5 mM phosphate buffer and a large excess (0.1 M) of inert NaClO4. Thiol solutions were made in 95% ethanol using >99% purity mercaptoundecylamine (AUT from Asemblon, Inc., Redmond, WA) and g98.5% purity 1-octanethiol (OT). The polycrystalline gold bead working electrodes were formed by melting 0.5 mm diameter gold wire (99.995%, Alfa Aesar). The approximate surface area was 0.2 cm2. The procedure for electrochemically polishing and conditioning these beads is provided elsewhere.46 Before each experiment, the gold beads were cleaned in freshly prepared Piranha solution (3:1 H2SO4/H2O2), rinsed with Millipore water, and then flame-annealed to provide a contaminantfree electrode surface. SAM Formation. After the cleaning procedure, the gold bead electrode was rinsed with ethanol and then immersed in the thiol incubating solution. This solution consisted of either 0.1 mM AUT and 1.5 mM OT (mixed composition SAMs) or 1.5 mM AUT (single-component SAMs). To ensure the formation of the amine-terminated monolayer, ammonium hydroxide was added to the incubating solution.47 After 1 h of incubation, the SAM was rinsed with ethanol and then immediately placed in a freshly prepared 5 mM quinone ethanolic solution for 3 min at 50 °C. The electrode was removed and then rinsed, first, with ethanol and then with Millipore water. The modified working electrode was dried with a stream of argon before being placed in the electrochemical cell. Electrochemical Measurements. All the glassware was cleaned in hot acid (1:3 HNO3/H2SO4), thoroughly rinsed with Millipore water, and then soaked overnight. Cyclic voltammetry and double-step chronocoulometry were performed using a HEKA PG590 potentiostat (HEKA, Mahone Bay, NS, Canada). Data were collected using a multifunction DAQ card (PCI 6251 M Series, National Instruments Austin, Texas) and in-house software written in the LabVIEW environment. Digitized chronoamperograms were numerically integrated. The resulting charge versus time plots were fit to a linear function at long times and extrapolated back to zero time to account for any current offsets. All of the experiments used a coiled loop of gold wire as the counter electrode and a Ag/AgCl electrode as reference electrode. Electrolytes were degassed with argon before the introduction of the SAM-quinone modified working
Figure 1. Cyclic voltammograms in pH 4.5 phosphate buffer electrolyte recorded at 5 mV/s for benzoquinone-derivatized selfassembled monolayers formed from ethanolic solutions of 1.5 mM AUT (----) and a 15:1 mixture (1.5 mM total thiol concentration) of OT/ AUT (s).
electrode, and a positive overpressure of argon was maintained above the electrolyte during the measurements that were performed at room temperature (21 ( 1 °C). Experiments were done in triplicate to afford a measure of experimental uncertainty. 3. Results and Discussion General Cyclic Voltammetry Features. When SAMs were formed from 1.5 mM ethanolic solutions of 11-aminoundecylthiol (AUT) and then derivatized with benzoquinone, the resulting 5 mV/s CVs in pH 4.5 phosphate buffer electrolyte were characteristic of the dotted line shown in Figure 1. This voltammogram reveals two redox couples whose apparent formal potentials are separated by ∼0.3 V, consistent with previous reports on similar benzoquinone-derivatized SAMs. Nova´k and co-workers have been able to demonstrate that post facto derivatization of amine-terminated SAMs with benzoquinone results in multiple binding motifs between the redox center and the functionalized SAM.32,33 Nucleophilic attack of the amine at the 2 position of benzoquinone produces singly bound redox centers. However, a terminal amine of an adjacent SAM molecule can further attack at the 5 position, resulting in disubstitution and a molecule that is doubly tethered to the electrode surface. As explained by Nova´k and co-workers, each addition reaction shifts the formal potential cathodically. For completeness, we note that in some instances a minor third redox couple was observed at more positive potentials and with much smaller peak currents (viz., the small redox signal observed at +0.2 V in the solid line in Figure 1). This third redox pair has been attributed to quinones noncovalently attached to the monosubstituted species.32 Using Nova´k’s interpretation, our CV for the one-component SAM reveals a much higher loading of disubstituted compared with singly bound benzoquinone. Even at very slow scan rates (1 mV/s), all four peaks displayed peak half-widths well in excess of the theoretical 45 mV expected for a 2e- surface redox couple.48 This nonideal behavior is expected for surfaces with strongly interacting redox centers resulting from high loading. The heterogeneity of such systems complicates the interpretation of kinetic studies where the measured parameters are averages of a wide distribution of microenvironments, each with their own characteristic standard heterogeneous rate constant. As strongly interacting redox centers mask many of the features we wished to explore, we attempted to dilute the benzoquinone surface density by using a SAM formed from a
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15:1 mixture (1.5 mM total thiol concentration) of octanethiol (OT) and AUT. This was motivated by the improved electrochemical reversibility reported by Kaifer et al. upon the dilution of anthraquinone SAMs with alkanethiols.49 Octanethiol was chosen as the diluting thiol to ensure that the quinone centers extended beyond the hydrophobic core of the monolayer.42 A representative 5 mV/s CV in pH 4.5 buffer is shown as the solid line in Figure 1. Integration of the CVs reveals a lower loading of quinone on the mixed SAM compared with the singlecomponent AUT monolayers. This difference is smaller than the expected factor of 1/15. Early studies of thiol SAM formation from the Whitesides group demonstrated that the equilibrium mole fraction of thiols assembled on a surface is not the same as the mole fraction in the incubating solution when using binary thiol solutions.50,51 It is quite likely that the fraction of the longer-chained AUT in our monolayers is significantly larger than 0.07%. In addition, in the AUT-only monolayer, the amine termini are presumably extensively hydrogen-bound and sterically hindered, which would cause the ratio of covalently attached benzoquinones to amine groups to be lower than unity. In comparison, the unconstrained amines on the mixed SAM surface are expected to be more reactive and provide larger benzoquinone-to-AUT ratios. Unlike the AUT-benzoquinone SAM, the mixed monolayer provides wellresolved voltammetric features. The CV consists of a strong, quasi-reversible redox couple at ∼0 V and a much weaker couple at approximately -0.3 V. These E0′ values are close to theapparentformalpotentialsobservedfortheAUT-benzoquinone SAM and indicate that the AUT molecules are sufficiently mixed in the OT SAM such that the overwhelming majority of benzoquinones are singly bound to the surface. Such systems were quite stable, although we did observe a slow attenuation of peak heights during prolonged scanning experiments. In all our subsequent analyses, we used 15:1 (OT/AUT) mixtures to build our base monolayers and only characterized the principal electrochemical feature corresponding to the monosubstituted quinones. Even though Figure 1 indicates that a noninteracting, homogeneous system can be produced from the mixed SAM, it also reveals extremely sluggish PCET kinetics. Scanning the potential as slow as 1 mV/s still resulted in full widths at halfmaximum (fwhm) between 50 and 75 mV, which is greater than the theoretical value of 45 mV for an ideal, Nernstian system but smaller than the fwhm values reported by Larsen and Gothelf for diluted mercapto-octylhydroquinone monolayers.42 Chronocoulometric Studies at Equilibrium Conditions. The very slow kinetics of the benzoquinone reduction prevents us from measuring the system under equilibrium conditions using potential sweep techniques. To obtain Nernstian behavior, we employed the following double-step chronocoulometric method. The working electrode was initially biased at a potential, Erest, of ∼0.2 to 0.3 V negative of the formal potential of the monosubstituted redox couple. Care was taken to ensure that this potential was positive of the residual signal arising from disubstituted benzoquinones. At Erest, all monosubstituted redox centers will be in the fully reduced form. The potential was then stepped to a more positive potential, Evar, and held for 2 min, which is long enough that the fraction of oxidation mandated by the Nernst equation is achieved. The potential was then stepped back to Erest, and the resulting current transient was measured for 250 ms. Erest is a sufficiently negative enough overpotential that all oxidized species formed in the forward step can easily be reduced within the 250 ms window even with standard heterogeneous rate constants on the order of 10-2 s-1. The transient was then numerically integrated to provide the
Zhang et al. difference in the total charge, ∆Q, between Erest and Evar. This procedure was repeated for increasingly positive values of Evar in 15 mV steps. The total charge measured consists of contributions from both Faradaic and capacitive processes
∆Q ) ∆qF + ∆qC
(1)
with only the former carrying analytical information concerning the benzoquinone speciation. Figure 2a provides the results of double-step experiments for both pH 1.9 and pH 5.5 phosphate buffer electrolytes. At potentials well-removed from the formal potentials, ∆Q varies linearly with potential with a constant slope. This indicates that the interfacial capacitance associated with the hydrophobic core of the SAM is potential-independent, in agreement with the CV results. Furthermore, the linearity allows us to accurately correct the ∆Q values to give ∆qF, as illustrated in Figure 2a. Figure 2b shows the fractional amount of oxidized redox centers as a function of potential, which can be determined from the measured charge as follows
θ)
Γox (∆qF)E ) Γtot (∆qF)tot
(2)
Γox is the surface concentration of the oxidized form of the monosubstituted benzoquinone. Γtot is the total surface concentration of monosubsituted redox centers, which is proportional to the total charge (∆qF)tot associated with the complete conversion of the reduced species to its oxidized form. Numerical differentiation of Figure 2b yields a plot of dθ/dE versus E, which is the equivalent of a linear sweep voltammogram under Nernstian conditions. Figure 2c provides such plots for the two pHs. The full width at half-maxima for pH 1.9 and pH 5.5 were measured to be 50 and 51 mV, respectively, proving that both redox waves correspond to nearly ideal, but very slow, two-electron processes. Influence of pH on Voltammetry. Figure 3 provides background-corrected CVs recorded at 1 mV/s for various pHs at constant ionic strength. Each CV was recorded on a freshly prepared electrode, and the CVs have been offset on the current axis for clarity. As it is impossible to ensure the same benzoquinone loading for every electrode preparation, the integrated peak areas are not identical in each CV. The voltammograms exhibit a high degree of asymmetry (most noticeable at pH 7.4), which is related to PCET kinetics (vide infra). For all pHs, the asymmetry becomes more pronounced at higher sweep rates. Figure 3 demonstrates that the redox couple shifts anodically with decreasing pH. Voltammograms recorded in electrolytes either more acidic than pH 1 or more basic than pH 9 produced very poor quality CVs with unstable peaks. We believe that this may be caused either by decomposition of the benzoquinone or by hydrolysis of the carbon-nitrogen bond. In discussing the pH dependence of PCET thermodynamics in our system, it is convenient to refer to the nine-membered square shown in Scheme 1. For convenience, we have used the same labeling convention as Laviron.5 In our particular case, M would represent the fully oxidized, unprotonated benzoquinone and V the fully reduced, dihydrobenzoquinone. Laviron2,5 (and more recently Finklea7) have theoretically treated various permutations and shown that, for both 1e1H PCET (e.g., the four-membered square scheme defined by the subset OPQR in Scheme 1) and 2e2H PCET (the full, nine-membered scheme), the observed formal potential should have a -60 mV pH-1
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Figure 3. Baseline-corrected 1 mV/s CVs of benzoquinone-derivatized 15:1 (OT/AUT) mixed SAMs in 0.1 M NaClO4 + 5 mM phosphate buffer electrolytes, the pH of which is indicated next to each curve. The CVs have been displaced along the ordinate axis for clarity.
SCHEME 1: 3 × 3 Square for a Two-Electron, Two-Proton Redox Couple
Figure 2. Results of chronocoulometry measurements for benzoquinone-derivatized 15:1 (OT/AUT) mixed SAMs in sodium phosphate buffer electrolytes. Filled squares (9) correspond to experiments performed at pH 5.5, and open circles (O) correspond to pH 1.9. (a) Relative charge as a function of stepped potential. (b) Fractional composition of oxidized species as determined from the charge measurements. (c) Numerical differentiation of panel b (points) and resulting Gaussian fits (solid lines). fwhm of fits are 50 mV for pH 5.5 and 51 mV for pH 1.9.
dependence. In the case of 1e1H PCET, this slope holds for electrolytes where the pKa1 < pH < pKa2 but has much more complicated limits for the 2e2H case. PCET involving unequal numbers of protons and electrons (1e2H and 2e1H) have expected slopes of -120 and -30 mV pH-1, respectively. From the results shown in Figure 3, we were able to extract the apparent formal potentials (estimated from the average anodic and cathodic peak positions in 1 mV/s scans) that are plotted versus pH in Figure 4. For 4.5 < pH < 8.5, the observed slope is almost exactly -60 mV pH-1, congruent with either a 1e1H or a 2e2H PCET. As the chronocoulometric experiments described above clearly reveal a two-electron process, we can exclude the possibility of 1e1H charge transfer and assert that, in this pH range, the observed electrochemical signals arise from the 2e2H conversion between the dihydrobenzoquinone and the benzoquinone. However, Figure 4 also shows that, in acidic buffer solutions, the 2e2H transfer may not be operative. Below pH 4.5, a slope of -88 mV pH-1 is observed, which is inconsistent with any of the permutations possible in the ninemembered square scheme. March et al. have also reported a nonconforming slope of -50 mV for the expected 2e2H transfer
in a 5-hydroxy-naphthoquinone monolayer that they believe is related to the presence of the 5-hydroxy group.52 From a simple consideration of the Nernst equation, a -90 mV pH-1 slope would arise from a process involving an overall transfer of two electrons and three protons. Our data in Figure 2 confirm that the observed reaction still involves two electrons at low pH, and therefore, we propose that a third protonation during the reduction of the benzoquinone to the dihydrobenzoquinone must be occurring. We believe that this third H+ transfer involves protonation of the secondary amine that bonds the AUT with the benzoquinone. This mechanism would be consistent with Abhayawardhana and Sutherland’s proposed electric-field-driven protonation/deprotonation of the aryl amine for an aminoanthraquinone self-assembled monolayer.31 Such field-assisted proton-transfer processes have been observed in redox inactive SAMs containing carboxylic acid functional groups.53–56 We are currently using electrochemical infrared spectroscopy to further investigate this proposed third protonation event. Influence of pH on PCET Kinetics. Using classical electrontransfer theory (potential-independent charge-transfer coefficients), Laviron derived the pH dependence on the apparent rate constant for the 2e2H example of PCET.5 Finklea subsequently modified Laviron’s approach to account for Marcus density of states theory by introducing potential-dependent transfer coefficients.7 The subtleties of the two approaches are difficult to resolve in 2e2H processes, as the apparent transfer coefficient, Rapp, becomes potential-dependent for both models. In either theory, simulated Rapp(η) plots tend to be asymmetric about the apparent formal potential and the shape of the asymmetry changes with changing pH. Finklea and Haddox observed this behavior in the CVs of galvinol/galvinoxyl SAMs44,45 and were able to simulate the experimental data using the potential-dependent charge-transfer coefficient model for 1e1H PCET. We qualitatively see the same affect in Figure 3 with 1 mV/s CVs at pH 7.4 and pH 4.4, showing the highest degree of asymmetry (the effect is even more pronounced at higher scan rates). Theoretical treatments for both classical and
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Zhang et al.
Epa ) E0′ -
Figure 4. Formal potential of benzoquinone-derivatized 15:1 (OT/ AUT) mixed SAMs in 0.1 M NaClO4 + 5 mM phosphate buffer electrolytes as a function of pH. The formal potential was estimated as the midpoint of the cathodic and anodic peak potentials in very slow scan (1 mV/s) CVs.
(3)
(1 - R)nFυa RT ln (1 - R)nF RTkapp
(4)
where Epc and Epa are the potentials of the cathodic and anodic peaks, n is the number of electrons transferred, and υa and υc are the critical scan rates. These last two parameters are obtained for both the anodic and the cathodic branches by plotting Ep E0′ versus ln(υ) and extrapolating the linear portion (which occurs when |E - Ep| > 100 mV/n) back to the x-axis intercept. The slopes of these lines also provide the transfer coefficients values that can be used to determine the apparent heterogeneous rate constants. Ideally, the rate constants and transfer coefficients obtained from the two branches of the CVs are self-consistent.
RnFυc RT
(5)
(1 - R)nFυa RT
(6)
kapp,c )
kapp,a )
Figure 5. Potential separation between the cathodic and anodic peaks as determined from 1 mV/s voltammograms.
density of states-dependent electron transfer also predict a pH dependence on the apparent standard heterogeneous rate constant. We will not reproduce the derivations but simply indicate that the relationship between log(ks,app) versus pH yields a “W” plot, consisting of two minima, centered around a local maximum. At pH extremes, the rate constant increases linearly as the pH moves further from neutrality, provided the 2e2H mechanism remains operative. The end result of Finklea’s adaptation of Laviron’s work is a modified “W” plot with an order of magnitude deeper local maximum and minima slightly shifted along the pH axis. Even without attempting to evaluate the apparent rate constants, we can see evidence of the predicted pH dependence by plotting the peak separation between the cathodic and anodic peaks as a function of pH, as shown in Figure 5. We now discuss several possible approaches to obtain ks,app. Two methods are based on voltammetric measurements, and one is a potential-step technique. Creager has also developed an AC voltammetry method,57 but we were unable to employ it for this study because our very slow ks,app values require measurements at inaccessibly low frequencies. Perhaps the most commonly used approach to obtain heterogeneous rate constants for electron transfer in redox-active SAMs is the Laviron approach. In this method, the peak separation is measured as a function of scan rate. Kinetic parameters may be obtained from eqs 3 and 458
( ) ( )
RnFυc RT ln RnF RTkapp
Epc ) E0′ -
Inherent in this approach is the assumption of a constant value of R at all peak potentials that move further from the formal potential with increasing scan rate. This assumption is not valid for PCET where Rapp is potential-dependent. We evidence this in Figure 6 where Laviron plots are presented for pH 3 and pH 7. At the higher pH, the anodic branch is quite linear but the cathodic branch fits to a second-order polynomial. The situation is reversed for the lower pH. Any attempt to fit the data to a linear function should result in serious error when extracting values for Rapp and kapp,s, and for this reason, we avoided using this approach. The slow kinetics of our system allowed us to employ a chronocoulometric method. If one writes the 2e2H reduction in terms of generalized oxidized and reduced species kc,app
Ox + ne- {\} Red
(7)
ka,app
Figure 6. Laviron plots for the anodic and cathodic branches of voltammograms recorded in pH 7 (9) and pH 3 (O) phosphate buffer electrolytes. Solid lines are either linear or second-order polynomial fits for data where |E - E0′| > 0.05 V.
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Figure 7. Kinetic results associated with a potential step from η ) 0 V to η ) -0.3 V as a function of hold time at the formal potential. Main plot: the left axis is the measured Faradaic charge for each transient, and the right axis is the corresponding fractional amount of the oxidized species. The inset shows the results of charge data linearization that can be used to extract the apparent heterogeneous rate constant, ks,app.
then the solution to the resulting differential equation for a potential step from θ ) 0 (Erest) to 0 < θ < 1 (Estep) can be written in terms of the fractional amount of oxidized redox centers and the apparent rate constants of the cathodic (kc,app) and anodic (ka,app) half reactions at Estep.
θ(t) )
ka,app [1 - e-kt] k
(8)
where k ) ka,app + kc,app. If Estep is chosen to be the formal potential, then ka,app ) kc,app ) ks,app and the fractional concentration of oxidized centers will reach θ ) 1/2 after sufficient time. By varying the duration, the potential is held at E0′, one can determine θ(t) by integrating the current transient arising from stepping the potential back to Erest. After correcting the total charge for background contributions (as described above), the measured charge [∆qF(t)]E is converted to θ(t), as per eq 2. The results of a typical experiment are shown in Figure 7, which shows that θ ) 1/2 is achieved after step durations on the order of 2 min at pH 7.5. Extraction of ks,app is achieved via eq 9 and linear regression (see the inset of Figure 7).
ln[1 - 2θ(t)] ) -2tks,app
(9)
As an alternative to the double-step chronocoulometric method, Finklea has described an approach to obtain ks,app based on cyclic voltammetry.44,45 This method involves isolating the Faradaic current from the double-layer charging currents and integrating the corrected voltammogram. At the apparent formal potential, the apparent standard rate constant and the Faradaic current, iF,η)0, are related as follows
ks,app )
iF,η)0 qF,tot[1 - 2xη)0]
(10)
In eq 10, xη)0 is the fraction of benzoquinone species reduced in the cathodic sweep from the positive limit to the formal potential. Equivalently, xη)0 can be computed for an anodic scan by integrating this linear sweep from the negative potential limit to the formal potential. Thus, from a single voltammogram, ks,app
Figure 8. Semilogarithmic plot of the apparent heterogeneous rate constant versus pH for the aminobenzoquinone monolayer as a function of pH. Data were obtained from cyclic voltammetry (solid squares) and double-step chronocoulometry (open squares).
can be calculated for both the anodic and the cathodic half cycles and an average value reported. In theory, the measurement can be performed for all scan rates as long as the peak separation is sufficiently large that the denominator does not approach zero but small enough that the Faradaic current at η ) 0 is sufficiently above background. The latter issue was the most demanding for our system and effectively prevented us from using all but the slowest scan rates (1 mV/s) for our kinetic analysis. In Figure 8, we show the results of our kinetic measurements as a function of pH using the two approaches detailed above. The measured apparent standard rate constants range between 0.1 and 0.01 s-1. Our values are larger by about an order of magnitude than the values reported earlier for similar length quinone monolayer systems.20,35 We note that, in previous studies, the kinetic information was extracted using the Laviron formalism which does not account for the potential dependence of Rapp. Quantitatively, the curve obtained from the potential step experiments is consistently larger than the data obtained from voltammetry. We believe that this systematic discrepancy arises from uncertainties in determining and then applying the true formal potential during the double step experiments. For a two-electron process, even small excursions from zero overpotential lead to large differences between the rate constants and the standard rate constant. Even an offset as little as ( 4 mV can lead to a nearly 20% change in the anodic and cathodic rate constants. Any difference between the actual stepped potential and the true formal potential, regardless of sign, results in an overestimate of ks,app when the double-step methodology described above is applied. Thus, the open data points in Figure 8 can be considered as the upper bound of apparent standard rate constants. Ideally, one could repeat the double-step measurement for a series of applied potentials near the perceived formal potential to better identify ks,app. This approach would also allow one to determine the apparent transfer coefficient as a function of potential. Unfortunately, our preliminary attempts to make these measurements failed due to the slow loss of quinone electroactivity with prolonged exposure to the electrolyte solution. Uncertainties in the formal potential will also cause errors when determining the rate constant from voltammetry. For an anodic scan, positive errors in E0′ will lead to overestimates of ks,app and negative errors in E0′ will give underestimates. Due to the asymmetry of the voltammograms, these errors will not be canceled out by averaging the results of the anodic and cathodic scans. Qualitatively, the two techniques provide very similar results, particularly in that they both show a clear minimum at pH ∼
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7.5. They also demonstrate that the apparent rate constant varies by nearly an order of magnitude in the range of 1 < pH < 9. This is consistent with the spread of apparent rate constants calculated by Finklea over a similar pH range for simulation conditions pertinent to benzoquinone.7 Although the decomposition of the aminobenzoquinone prevents us from measuring ks,app at pHs higher than 8.5, Figure 8 implies increasingly facile kinetics with increasing pH. However, it is less clear if this trend is mirrored at low pH. Once again, the system’s instability at pH extremes prevents us from extending our studies over a wider range of electrolyte acidities. Whereas the primary minimum at pH ∼ 7.5 is well-defined, the secondary minimum is less pronounced than expected. Apparent rate constants extracted from the charge and CV measurements reveal a shallow minimum at pH ∼ 4.5. The position and width of the weaker minimum are very consistent with the peak separation data shown in Figure 5. We believe that, although only weakly expressed, there is a secondary minimum in the rate constant versus pH plot centered at pH 4.5 and that the overall shape and appearance of the curves in Figure 8 provides a distorted “W” plot. The difference between our experimental curves and Finklea’s simulations7 may arise from shifts and broadening of the bulk pKa values of quinone upon confinement to the SAM/ electrolyte interface. It is also unclear to us what role the third protonation at low pH plays in the PCET kinetics. In the theoretical description of proton-coupled electron transfer, the (de)protonations are believed to occur much more rapidly than the electron transfers and the acid-base chemistry is considered to be at equilibrium. Therefore, one might expect the kinetics of the 3H2e system to be equivalent to the case of 2e2H transfer. That assumes protonation of the secondary amine does not provide new electron-transfer pathways (i.e., no difference in the formal potentials of the reduction of the ammoniumbenzoquinone and the aminobenzoquinone). 4. Summary and Conclusions We have examined the proton-coupled electron-transfer reactions of a benzoquinone-modified self-assembled layer. Our preparation technique has allowed us to obtain nearly ideal electroactive monolayers with little to no interaction between redox centers. The ideality of our system greatly reduces the heterogeneity of our measured thermodynamic and kinetic parameters. The dependence of the apparent formal potential on pH displays two linear regions. Between 4.5 < pH < 8.5, the observed slope is -58 mV, consistent with 2e2H transfer. In more acidic electrolytes, the absolute value of the slope increases to 88 mV. In low pH electrolytes, the number of electrons transferred is still two, implying that a third proton transfer occurs during the PCET (2e3H). We speculate that an electricfield-driven protonation of the alkylaminobenzoquinone is responsible for the third proton transfer. Efforts are now underway in our laboratory to provide spectroscopic evidence of the third protonation. Our observations of asymmetric cyclic voltammograms and variable peak separations indicate that the apparent standard rate constant is also pH-dependent. We have outlined two means to determine ks,app that do not require the assumption of a potentialindependent apparent transfer coefficient. The two approaches are in reasonable qualitative and quantitative agreement and provide a distorted version of the W plot predicted from theory for stepwise 2e2H proton-coupled electron transfer. We note that the position of the less-pronounced minimum is close to the transition pH where the system switches from 2e2H to 2e3H transfer. Although a theoretical framework for 2e3H transfer
Zhang et al. has not yet been developed, one might expect such a system to yield a third minimum due to proton transfer on the amine center rather than protonation of the quinone. It is possible that the minimum at pH 4.5 in Figure 8 is unique to the aminobenzoquinone system and the true second minimum predicted for 2e2H transfer may occur at higher, and unfortunately inaccessible, solution pH values. Interpreting Figure 8 as a distorted W plot inherently assumes that PCET in this surface-confined benzoquinone system follows a stepwise mechanism. The validity of this presumption is perhaps less certain after Finklea’s studies of 1e1H transfer in galvinol SAMs44,45 and monolayers of osmium aquo complexes.59,60 Neither of these systems adequately displayed the required pH dependence on either the standard rate constant or the transfer coefficient and led Finklea to speculate that a concerted, rather than a stepwise, PCET mechanism is operative. At this juncture, we cannot determine if the minor deviations from the expected log(ks,app) dependence on pH can be explained as perturbations to stepwise PCET arising from the particulars of our system or whether they are better explained by a concerted PCET mechanism. To address this question, efforts are now underway in our laboratory to extract the potential dependence of Rapp and to investigate kinetic isotope effects.60,61 Acknowledgment. This work was supported by a grant from the Natural Sciences and Engineering Research Council of Canada (NSERC). References and Notes (1) Laviron, E. J. Electroanal. Chem. 1980, 109 (1-3), 57–67. (2) Laviron, E. J. Electroanal. Chem. 1981, 124 (1-2), 1–7. (3) Laviron, E. J. Electroanal. Chem. 1981, 124 (1-2), 9–17. (4) Laviron, E. J. Electroanal. Chem. 1983, 146 (1), 1–13. (5) Laviron, E. J. Electroanal. Chem. 1983, 146 (1), 15–36. (6) Laviron, E. J. Electroanal. Chem. 1984, 164 (2), 213–227. (7) Finklea, H. O. J. Phys. Chem. B 2001, 105 (37), 8685. (8) Baxendale, J. H.; Hardy, H. R. Trans. Faraday Soc. 1953, 49, 1140– 1144. (9) Chambers, J. Q. Electrochemistry of Quinones. In The Chemistry of Quinoid Compounds; Patai, S., Ed.; Wiley: New York, 1974. (10) Bailey, S. I.; Ritchie, I. M. Electrochim. Acta 1985, 30 (1), 3–12. (11) Forster, R. J.; O’Kelly, J. P. J. Electroanal. Chem. 2001, 498 (12), 127–135. (12) Mu¨ller, O. H. J. Am. Chem. Soc. 2002, 62 (9), 2434–2441. (13) Quan, M.; Sanchez, D.; Wasylkiw, M. F.; Smith, D. K. J. Am. Chem. Soc. 2007, 129 (42), 12847. (14) Shim, Y. B.; Park, S. M. J. Electroanal. Chem. 1997, 425 (1-2), 201–207. (15) Deakin, M. R.; Wightman, R. M. J. Electroanal. Chem. 1986, 206 (1-2), 167–177. (16) Costentin, C.; Robert, M.; Saveant, J. M. J. Am. Chem. Soc. 2006, 128 (27), 8726–8727. (17) Costentin, C.; Robert, M.; Saveant, J. M. J. Am. Chem. Soc. 2007, 129 (18), 5870–5879. (18) Costentin, C. Chem. ReV. 2008, 108 (7), 2145–2179. (19) Hammes-Schiffer, S.; Soudackov, A. V. J. Phys. Chem. B 2008, 112 (45), 14108–14123. (20) Trammell, S. A.; Seferos, D. S.; Moore, M.; Lowy, D. A.; Bazan, G. C.; Kushmerick, J. G.; Lebedev, N. Langmuir 2006, 23 (2), 942. (21) Trammell, S. A.; Lowy, D. A.; Seferos, D. S.; Moore, M.; Bazan, G. C.; Lebedev, N. J. Electroanal. Chem. 2007, 606 (1), 33–38. (22) Trammell, S. A.; Moore, M.; Lowy, D.; Lebedev, N. J. Am. Chem. Soc. 2008, 130 (16), 5579. (23) Trammell, S. A.; Moore, M.; Schull, T. L.; Lebedev, N. J. Electroanal. Chem. 2009, 628 (1-2), 125–133. (24) Kwon, Y.; Mrksich, M. J. Am. Chem. Soc. 2002, 124 (5), 806. (25) Yousaf, M. N.; Mrksich, M. J. Am. Chem. Soc. 1999, 121 (17), 4286. (26) Muhammad, N. Y.; Eugene, W. L. C.; Milan, M. Angew. Chem., Int. Ed. 2000, 39 (11), 1943–1946. (27) Katz, E.; Schmidt, H. L. J. Electroanal. Chem. 1993, 360 (1-2), 337–342. (28) Lukkari, J.; Kleemola, K.; Meretoja, M.; Kankare, J. Chem. Commun. 1997, 1099–1100.
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