J. Phys. Chem. C 2009, 113, 4915–4924
4915
Heterogeneous Proton-Coupled Electron Transfer of an Aminoanthraquinone Self-Assembled Monolayer Anusha D. Abhayawardhana and Todd C. Sutherland* Department of Chemistry, UniVersity of Calgary, 2500 UniVersity DriVe NW, Calgary AB, Canada T2N 1N4 ReceiVed: August 14, 2008; ReVised Manuscript ReceiVed: January 13, 2009
Electrochemical properties of self-assembled monolayers (SAMs) of a 1-aminoanthraquinone derivative (1aminoAQ) were studied using electrochemical impedance spectroscopy (EIS) and cyclic voltammetry. The results of EIS studies at different temperatures and applied potentials established that SAMs of 1-ammoniumAQ undergo proton-coupled electron transfer (PCET) in 0.1 M H2SO4 with a rate constant of 7.4 s-1. Variabletemperature studies of the charge-transfer reaction gave a high activation barrier of 69 kJ · mol-1 for the 1-ammoniumAQ redox reaction, which is attributed to cleavage of a strong intramolecular hydrogen bond; this conclusion is supported by NMR and FTIR spectroscopic analyses, as well as DFT calculations. The 1-ammoniumAQ redox reactions were found to have a reorganizational energy term of 2.7 eV, consistent with large solvent reorganization and substantial nuclear rearrangement upon reduction. Introduction 1
Electron-transfer and proton-coupled electron-transfer reactions are critical steps in both photosynthesis and cellular respiration, especially considering that quinone redox reactions occur in every living organism. The redox chemistry of quinone reduction to dihydroquinone is a well-documented two-electron, two-proton (2e- + 2H+) reduction, and the mechanism has been studied and debated for decades.2-8 The redox mechanism has been shown to change as a function of both solvent and pH. In dry polar solvents, the cyclic voltammagram (CV) shows two distinct one-electron transfer reactions.4,9,10 In aqueous solutions, the 2e- + 2H+ redox reaction changes, and the CV peaks often coalesce and manifest as one redox wave.4,11 Recent work by Costentin, Robert, and Save´ant2,12,13 suggests a concerted mechanism for electron and proton transfer that could be operating in several natural processes, such as those catalyzed by the enzymes cytochrome c oxidase,14 ribonucleotide reductase,15 and superoxide dismutase.16 Under basic conditions (pH > 10),17,18 quinones show reversible redox waves by cyclic voltammetry, but the E0 values remain constant, which is an effect of the pH exceeding the pKa of the phenolic protons, consistent with a mechanism that involves only two electrons. In addition, at high pH, the electron-transfer rate is higher than that under less basic conditions.19 At physiological pHs (between pH 5 and 8), the mechanism has been shown to change from the sequence H+, e-, H+, e- below pH 5 to the sequence e-, H+, e-, H+ above pH 6.20-22 Studies of the effect of pH on E0 have shown the slope at pH < 10 to be 60 mV · pH-1, consistent with Nerstian behavior of a 2e- + 2H+ redox process. Under neutral pH conditions, the electron transfer is clearly coupled to proton transfer, and the electron-transfer rate is dramatically reduced compared to basic conditions. Quinones under acidic conditions are expected to couple proton-transfer reactions with their redox reaction. The stepwise mechanism is unclear, but evidence points toward a 2e- + 2H+ process,20,23 similar to that occurring under neutral conditions. The kinetics of charge transfer is slower under acidic conditions, and the E0 values * To whom correspondence should be addressed. E-mail: todd.sutherland@ ucalgary.ca.
continue the trend of shifting toward anodic potentials at lower pHs. The quinone redox chemistry is further complicated by whether the redox chemistry is analyzed under buffered or unbuffered conditions.4 In addition, the role that hydrogenbonding solvents play in the intermediate has been shown to affect the redox chemistry.4,24 The majority of the quinone research work to date has been of redox probes in solution. Despite much research into the redox chemistries of quinones, areas remain that are less studied, especially the kinetics of charge transfer. Detailed mechanistic information is becoming available, but only a subset of the relevant studies have investigated the charge transfer of quinones at an interface. Interfacial charge transfer is critical to understanding charge transfer in biological systems because the quinone species is typically membrane-bound and the redox reaction occurs at the aqueous|biomembrane interface. The dynamics of charge transport, restricted motion, and solvation of a redox probe that lies at the interface between hydrophilic and hydrophobic layers are expected to be vastly different from those of a well-solvated diffusing redox probe. For example, a benzoquinone redox group at physiological pH has a half-potential (E1/2) of -0.122 V vs Ag|AgCl|KCl in an unbuffered solution,4 whereas the same benzoquinone redox group that has formed a self-assembled monolayer (SAM) using Au-thiol chemistry has a formal potential (E0) of +0.04 V vs Ag|AgCl|KCl.25 The difference in free energies between solution redox and surface-bound redox (∆∆G) is 31 kJ · mol-1, which demonstrates that the interface can influence the redox thermodynamics. This simplified comparison between solution and surface-bound benzoquinone redox chemistry does not take into account the alkyl thiol spacer group, (CH2)8, which can exert a large solvation difference and a small electronic effect on the benzoquinone. This contribution examines the redox reactions of aminoanthraquinone (aminoAQ) derivatized with an alkyl (C10H20) disulfide. AminoAQ was chosen because its redox chemistry closely parallels that of benzoquinone and the amine allows for greater synthetic versatility. The alkyl group was installed through a substitution reaction of the corresponding alkyl iodide with the aryl amine. Although synthetically useful, the amine is the most basic site on the molecule, which could be protonated
10.1021/jp807287p CCC: $40.75 2009 American Chemical Society Published on Web 02/27/2009
4916 J. Phys. Chem. C, Vol. 113, No. 12, 2009 in acidic media and complicate the proton-coupled redox chemistry. Alkyldisulfides have been shown to spontaneously form Au-thiolate monolayers when incubated with a clean Au surface.26 SAMs of numerous redox probes, including ferrocene,27-29 quinone,17,25,30-32 and anthraquinone,19,33-35 have been extensively studied because they form well-defined organic monolayers that can be used to understand charge transfer in mechanistic detail. The 10-carbon alkyl spacer group has been shown to have favorable van der Waals interactions, leading to a tightly packed monolayer configuration that limits the number of defect sites and allows for more reliable electrochemical characterization.26 An analysis of the thermodynamics and kinetics of charge transfer of an aminoAQ SAM on a gold surface at different unbuffered pHs is presented. Because little kinetic information is available, a thorough characterization of the charge transport at acidic pHs was required. This study examines the rate of charge transfer at different temperatures to derive activation barriers and provide mechanistic insight into the proton-coupled electron-transfer reaction by probing the aminoAQ SAM under different applied potentials using cyclic voltammetry and electrochemical impedance spectroscopy (EIS).
Abhayawardhana and Sutherland SCHEME 1: Synthesis of Anthraquinone-Terminated Decyldisulfidea
Experimental Section Materials and Measurements. 1-Aminoathraquinone (aminoAQ), 1,10-dibromodecane, potassium thioacetate, KCl, KI, tetrahydrofuran (THF), and sulfuric acid were all purchased from Sigma-Aldrich and used as received. All solvents were purchased from Sigma-Aldrich and deoxygenated and dried with an MBRAUN solvent purification system. Reactions were performed using oven-dried glassware under an atmosphere of nitrogen. Nuclear magnetic resonance (NMR) spectra were recorded on Bruker 400 Ultrashield and Avance DMX 300 instruments and referenced against CDCl3. UV-visible and Fourier transform infrared (FTIR) spectra were recorded on a Cary 5000 UV-vis-NIR spectrophotometer and a Varian FTS 7000 FTIR spectrophotometer, respectively. Au wire (99.99%), Pt mesh, and Ag wire (99.9%) were purchased from Alfa Aesar. Synthesis. Preparation of Anthraquinone DeriWatiWe. The formation of 1,10-diiododecane followed a general halogenexchange procedure36 of reacting excess KI with 1,10-dibromodecane in acetone at reflux for 4 h, and product identification was confirmed by comparison to literature 1H and 13C NMR data. Scheme 1 details the synthesis of 1-aminoanthraquinone (aminoAQ) disulfide (3), which was used in the monolayer studies. Synthesis of 1. This synthesis is based on a similar procedure by Caruso et al.37 1-Aminoanthraquinone (0.04 g, 0.39 mmol) and 1,10-diiododecane (0.614 g, 1.56 mmol) were refluxed in 3 mL of acetonitrile for 5 days. The solution was concentrated under reduced pressure and purified by silica gel chromatography [hexane/ethyl acetate (EtOAc), 10:1; rf ) 0.36] to give 1 (0.045 g, 24% yield) as a maroon powder. 1H NMR (300 MHz, CDCl3) δ: 9.72 (s, 1H), 8.35-8.14 (m, 2H), 7.74 (td, J ) 1.6, 7.5 Hz, 1H), 7.67 (td, J ) 1.5, 7.4 Hz, 1H), 7.62-7.47 (m, 2H), 7.03 (dd, J ) 1.4, 8.2 Hz, 1H), 3.30 (dd, J ) 6.9, 12.3 Hz, 2H), 3.16 (t, J ) 7.0 Hz, 2H), 1.88-1.66 (m, 5H), 1.60-1.12 (m, 15H). 13C NMR (75 MHz, CDCl3) δ: 184.93, 183.86, 151.80, 135.26, 135.05, 134.64, 133.91, 133.02, 132.85, 126.69, 126.62, 117.88, 115.50, 112.79, 42.97, 33.51, 30.47, 29.38, 29.32, 29.28, 29.08, 28.49, 27.14, 7.33. Anal. Calcd (%) for C24H28INO2 (489.12): C, 58.9; H, 5.77; N, 2.86. Found: C, 57.78; H, 6.44; N, 2.57. ESI-MS m/z: [M + Na]+ calcd, 512.11; found, 511.81; relative peak intensity, 16.7%.
a Conditions: (i) acetonitrile, reflux 5 days; (ii) K+CH3(CO)S-, acetonitrile, reflux 6 h; (iii) NaOH, acetone, 5 days at room temperature.
Synthesis of 2. This synthesis is based on a similar procedure by Caruso et al.37 Both 1 (0.024 g, 0.049 mmol) and potassium thioacetate (0.027 g, 0.24 mmol) were mixed in acetonitrile (24 mL) and then refluxed for 6 h under N2 atmosphere. The acetonitrile was distilled off under reduced pressure, and the residue was suspended in CH2Cl2 and filtered. The organic phase was concentrated under reduced pressure and purified by silica gel chromatography (hexane/EtOAc, 15:1; rf ) 0.19) to give 2 (0.016 g, 75% yield) as a red solid. 1H NMR (300 MHz, CDCl3) δ: 9.73 (s, 1H), 8.32-8.13 (m, 2H), 7.73 (td, J ) 1.6, 7.5 Hz, 1H), 7.68 (td, J ) 1.6, 7.4 Hz, 1H), 7.61-7.45 (m, 2H), 7.04 (dd, J ) 1.5, 8.2 Hz, 1H), 3.39-3.29 (m, 2H), 2.87-2.80 (t, J ) 7.4 Hz, 2H), 2.30 (s, 3H), 1.80-1.66 (m, 2H), 1.60-1.12 (m, 17H). 13C NMR (75 MHz, CDCl3) δ: 196.08, 184.94, 183.89, 151.82, 135.26, 135.06, 134.64, 133.91, 133.02, 132.85, 126.69, 126.63, 117.90, 115.50, 112.79, 42.98, 30.64, 29.47, 29.40, 29.37, 29.29, 29.12, 29.09, 29.06, 28.78, 27.14. Anal. Calcd (%) for C26H31NO3S (437.20): C, 71.36; H, 7.14; N, 3.20. Found: C, 70.85; H, 7.39; N, 2.84. ESI-MS m/z: [M + Na]+ calcd, 460.19; found, 459.97; relative peak intensity, 70%. Synthesis of 3. This synthesis is based on a similar procedure by Gu et al.38 Reagent-grade acetone (2 mL) and 3 M NaOH (2 mL) were placed in a flame-dried 25 mL flask that was then deoxygenated with N2 gas purge for 20 min at 0 °C. Under positive N2 pressure, 2 (0.1 g, 0.23 mmol) was added, and the mixture was stirred for 5 days at room temperature. HCl (1 M) was then added to neutralize the solution to pH 7, and the aqueous fraction was extracted three times with EtOAc. The combined organic extracts were washed with brine (10 mL) and dried over MgSO4, and the solvent was evaporated under reduced pressure. The remaining residue was purified by silica gel chromatography (hexane/THF, 1:1; rf ) 0.33) to give 3
Heterogeneous PCET of an Aminoanthraquinone SAM (0.086 g, 48% yield) as a maroon powder. 1H NMR (300 MHz, CDCl3) δ: 9.72 (s, 1H), 8.28 (dd, J ) 1.4, 7.6 Hz, 1H), 8.24 (dd, J ) 1.4, 7.5 Hz, 1H), 7.79-7.61 (m, 2H), 7.61-7.44 (m, 2H), 7.03 (dd, J ) 1.5, 8.2 Hz, 1H), 3.30 (dd, J ) 7.0, 12.2 Hz, 2H), 2.65 (t, J ) 7.4 Hz, 2H), 1.83-1.62 (m, 2H), 1.54-1.16 (m, 19H). 13C NMR (101 MHz, CDCl3) δ 184.94, 183.84, 151.79, 135.25, 135.06, 134.65, 133.90, 133.03, 132.85, 126.70, 126.62, 117.88, 115.52, 112.83, 43.01, 39.18, 29.43, 29.42, 29.32, 29.19, 29.10, 28.49, 27.16. Anal. Calcd (%) for C48H56N2O4S2 (789.37): C, 73.06; H, 7.15; N, 3.55. Found: C, 72.14; H, 7.59; N, 3.35. ESI-MS m/z: [M + H]+ calcd, 789.38; found, 789.04; relative peak intensity, 11%. Electrochemistry. Electrochemical studies were carried out in an all-glass, thermal-jacketed, three-electrode custom-built cell. A custom-built Ag|AgCl|KCl3M electrode was used as a reference electrode, which was isolated from the working solution through a Vycor tip. A Pt mesh was used as a counter electrode, and a gold ball was used as the working electrode. The working electrode was cleaned by being heated in a butane flame until nearly melted and then being submerged in 18 MΩ · cm Milli-Q water, repeated three times. Geometric area was measured at 0.345 cm2. The Pt mesh counter electrode was heated with a butane flame for 5 min prior to each measurement. H2SO4 (0.1 M) was used as the electrolyte. The cell was thermoregulated using a circulating water bath (Lauda, Ecoline Refrigerating Circulators RE-200 Series), and all temperatures are reported as the internal electrolyte solution temperature. Ar gas bubbling for 30 min preceded each experiment, and a continuous Ar blanket above the electrolyte was applied during the entire electrochemical experiments. Cyclic voltammetry experiments were conducted at several scan rates from -0.4 to 0.4 V vs Ag|AgCl|KCl3M using an analog triangle waveform generator (Autolab PG302, ScanGen module). iR compensation was manually applied to each cyclic voltammetry experiment using the high-frequency intercept of the Nyquist plot as the compensation value (typically between 4 and 7 Ω; uncorrected for electrode area). Electrochemical impedance spectroscopy (EIS) was carried with on the same instrument using applied dc potentials and sweeping frequencies from 100 kHz to 5 mHz with an ac excitation signal of 10 mV and collection of 10 data points per decade. A PC using FRA and GPES version 4.9 software controlled the electrochemical system and was used for data analysis. Monolayers were prepared on the Au ball electrode by first cleaning the electrode using the method described above and then incubating the Au ball electrode in a sealed vial containing 1 mM of the electroactive disulfide dissolved in dry THF for 24 h. The Au electrode was then rinsed in large amounts of THF followed by Milli-Q water. Results and Discussion The SAM of 3 on a gold film (Evaporated Metal Films, Ithaca, NY) was structurally characterized by photoelastic modulation infrared reflection absorption spectroscopy (PEMIRRAS), the results of which are included as Supporting Information (Figure S10). The surface IR spectrum exhibits strong absorptions in the alkyl CsH stretching region showing the following three modes: methylene antisymmetric CsH stretch (d-, 2928 cm-1), methylene symmetric CsH stretch (d+, 2859 cm-1), and methyl antisymmetric CsH stretch (r-, 2967 cm-1). The surface reflection IR measurements indicate that the monolayer is relatively disordered based on the d- peak position.39,40 A well-ordered, semicrystalline monolayer shows antisymmetric CsH stretching at 2918 cm-1, and higher-
J. Phys. Chem. C, Vol. 113, No. 12, 2009 4917
Figure 1. Example cyclic voltammagrams of a SAM of 3 at scan rates ranging from 500 mV · s-1 to 20 V · s-1 in 0.1 M H2SO4 at 7.3 °C. Potentials are referenced to Ag|AgCl|KCl3M.
frequency stretches have been attributed to a greater number of gauche defects in the alkyl chain linker.41 The alkyl disorder was not unexpected, as this is a room-temperature measurement with a shorter C10 alkyl chain. The amine NsH and AQ CdO stretches were not observed, most likely because of the weak absorptions of these transitions. A typical cyclic voltammagram (CV) showing a single redox couple is shown in Figure 1 for a SAM of 3 at 7.3 °C in 0.1 M H2SO4. The CV shows a quasireversible charge-transfer reaction consistent with previously reported 2e- + 2H+ charge transfer and a formal potential (E0) of -137 mV vs Ag|AgCl|KCl3M, determined from the average of the peak potentials. The redox potential is substantially different from that of benzoquinone under similar conditions (0.215 V vs Ag|AgCl|KCl3M),42 but it compares well with those of anthraquinone derivatives bearing electron-donating groups, such as thiols (-0.12 V vs Ag|AgCl|KCl3M).35 Note that the redox potential in ref 35 was extrapolated from basic conditions (0.1 M KOH) assuming 0.059 V per pH unit. The CV peaks are asymmetric, with the oxidation peak having a full-width-at-half-maximum (fwhm) value that varies from 81 to 184 mV and the cathodic peak having fwhm values that vary from 63 to 154 mV. A reversible, Nerstian charge-transfer reaction should have a fwhm of 90/n mV.43 The ranges of fwhm values are typical for anthraquinone monolayers, with the lower values corresponding to lower scan rates. Two trends are clear from the CV. First, the peak heights of both the oxidative and reductive reactions increase with scan rate, which is expected assuming that the anthraquinone undergoes a quasireversible charge-transfer reaction. The peak currents (anodic, ipa, and cathodic, ipc) are linear with scan rate (see Figure S1 of the Supporting Information), indicating that surface-bound charge-transfer reactions, as opposed to diffusion-controlled redox reactions,43 were occurring. The second trend in the CVs is that the anodic peak potential (Epa) and the cathodic peak potential (Epc) separate with increasing scan rate, which is typical for an electroactive monolayer that is under kinetic control.43 Further analysis of the CV was done by subtracting the background charging current using software provided in the GPES v4.9 package. The resulting background-subtracted CV was then integrated to give the total charge (Qtot). Taking into account the electrode area and assuming a two-electron reduction per molecule, the total surface coverage (Γtot) was calculated at 1.7 × 10-10 mol · cm-2, which is consistent with other reports
4918 J. Phys. Chem. C, Vol. 113, No. 12, 2009
Figure 2. Charge-transfer rate analysis of a SAM of 3 using the Laviron formalism.
of quinone SAMs on gold.32,35,42,44,45 An ideal SAM of 3, assuming a 12 Å × 3.35 Å cross-sectional area, yields a coverage of 4.2 × 10-10 mol · cm-2, which is comparable to the measured coverage considering that the surface IR results show gauche defects, resulting in less-than-ideal monolayer packing. Further analysis of the CVs yields the rates of charge transfer using Laviron’s formalism.46 Figure 2 shows the plot of Epa and Epc versus the natural logarithm of the scan rate. The kinetic analysis presented by Laviron requires that the peak-to-peak separation be greater than 200 mV/n, which, in the case of quinones, for which n ) 2, limits the linear fit region to points greater than )/-50 mV of E0. A linear best fit to data points that meet the peak separation criterion leads to determination
Abhayawardhana and Sutherland of the charge-transfer coefficient (R) and the apparent rate constant of charge transfer (kapp) by the intercept of the fit line with Ep - E0 ) 0, termed the critical scan rate. The Laviron formalism gives R values approaching 0.5 and apparent chargetransfer rates, kapp, of 25, 44, and 49 s-1 for temperatures of 7.3, 16.1, and 34.6 °C, respectively. The rate constants calculated using the Laviron approach are dependent on an accurate measurement of the critical scan rate, and small changes in slope of the linear fit can result in large changes in measured rate constants. The Supporting Information includes a CV at a scan rate of 2 mV · s-1 (24.8 °C) that resulted in an unchanged formal potential and a peak separation of 38 mV. Ideal Nerstian behavior predicts that a surface-bound redox probe should have zero peak separation,43 and deviations from Nerstian behavior are commonly attributed to defects, nonideal packing, rough surfaces, or lateral interactions of redox probes. Electrochemical impedance spectroscopy (EIS) is a powerful tool for investigating redox-active monolayers because it is nondestructive and it probes a large time domain (microseconds to 103 s). Applying the formal potential, determined from the CV, to the electrode forces SAMs of 3 to be at equilibrium. Furthermore, by perturbing this equilibrium with a small (10 mV) ac excitation signal at many frequencies allows detailed investigations into the kinetics of the charge-transfer reaction. Figure 3 is a Bode representation of the EIS data for SAMs of 3 at different temperatures. The more common Nyquist representation is included as Supporting Information. The Bode plots clearly show two peaks at each temperature examined with a characteristic frequency labeled in the valley of the two peaks. Often, a peak in the Bode plot can be modeled with a resistor-capacitor (RC) circuit that can be interpreted chemically as a charge-transfer reaction.47 However, several equivalent circuits are capable of producing a peak in the Bode plot, and
Figure 3. EIS Bode plots for SAMs of 3 at an applied potential of -0.14 V vs Ag|AgCl|KCl3M with 10 mV excitation signal and a frequency range from 100 kHz to 5 mHz. All symbols are experimental data, and lines are best fits to the data. Temperature: (A) 7.3, (B) 16.1, (C) 24.8, and (D) 34.6 °C.
Heterogeneous PCET of an Aminoanthraquinone SAM
J. Phys. Chem. C, Vol. 113, No. 12, 2009 4919 EIS methods examine the system at steady state with only small sinusoidal, or ac, perturbations; thus, the kinetic parameters are considered more reliable than kinetic values derived from CV methods.50,51 The relationship between exchange current (i0) and charge-transfer resistance (Rct) shown in eqs 1 and 2 allows the kinetics of charge transfer (kct) to be calculated from Rct shown by eq 3
RT nFRct
(1)
i0 ) nFkctΓ
(2)
i0 ) Figure 4. Three equivalent circuits used to fit impedance data.
judicious model selection is needed. Figure 3 indicates that two processes are occurring in two time domains. EIS is interpreted by curve fitting the data to equivalent circuit models using the complex nonlinear least-squares (CNLS) technique, which is included in the FRA software. Three equivalent circuit models that describe SAMs of 3 are shown in Figure 4. Circuit 4A is a circuit that describes a polarizing electrode47 with terms Rs and Cdl in series, where Rs is the solution resistance and Cdl represents the double-layer capacitance. Circuit 4A does not model a charge-transfer reaction. The capacitor is actually a constant-phase element (CPE) that more accurately models rough electrodes. Throughout this contribution, CPE will be treated as a capacitor because the power-law modifier is typically between 0.8 and 0.9, where 1.0 is an ideal capacitor. Using CPE as a capacitor has limitations that can lead to large errors in estimating capacitance values; thus, capacitance values throughout this contribution should be considered as relative measures rather than absolute. For a more detailed discussion on the use of CPE as a capacitor, see Zoltowski.48 Circuit 4B is a parallel RC circuit, indicating that the monolayer under investigation has the same double-layer capacitance and an additional resistive element, shown as Rd, that represents defect sites or ion penetration.49 Circuit 4C models a charge-transfer reaction. The same circuit elements as in circuits 4A and 4B are employed, in addition to the chargetransfer resistance (Rct) and monolayer capacitance (Cml). The lines in Figure 3 are the results of the best fits of the data to the circuit model shown in Figure 4C. Table 1 details the results of the fitting procedure into discrete circuit elements. The solution resistance, Rs, for each temperature remains constant at ∼2.5 Ω · cm-2, which is a reasonable value considering the high electrolyte concentration. Note that temperatures greater than 35 °C were attempted, but these temperatures led to monolayer instability as assessed by the loss of aminoAQ redox chemistry over the course of an experiment. The initial CVs were compared to CVs recorded after the EIS procedure to ensure monolayer integrity over the several-hours-long EIS data collection period. All data shown in this work fulfilled the thermal stability criterion of less than 5% difference in integrated charge between the CVs recorded before and after EIS. The double-layer capacitance at each temperature is consistent at 30 µF · cm-2 with similar power-law modifier values of 0.86, suggesting little structural change in the monolayer at different temperatures. A clear trend in the resistance to charge transfer (Rct) is observed (from 696 to 58 Ω · cm2) with temperature, indicating faster electron transfer with increased temperature. The monolayer capacitance (Cml) values are scattered at the various temperatures, and little structural information is available. The defect site resistance (Rd) or ion penetration resistance is very large at 326 kΩ · cm2 at 7 °C and 22 kΩ · cm2 at 35 °C, consistent with small defect sites or large barriers to partitioning ions into the hydrophobic monolayer.
kct )
RT n F ΓRct 2 2
(3)
where R, T, n, and F are the gas constant, temperature, number of electrons involved in the charge transfer, and Faraday’s constant, respectively. Surface coverage, Γ (mol · cm-2), was calculated from the CV as described previously. The rate of the charge transfer at each temperature is reported in Table 1, which shows a reasonable correlation with the rates obtained by CV. Although the rate constants calculated by EIS and CV are close, a more thorough method to evaluate electron-transfer rates was carried out following the protocol presented by Brevnov et al.52 derived from the faradaic admittance theory described by Los and Laviron53 and Laviron.54,55 Figure 5 shows the stepwise procedure to remove nonfaradiac processes from the admittance data. In this protocol, the faradaic impedance is extracted from other impedances (Rs and Cdl, as a constant phase element). Figure 5A is the total cell admittance including both real and imaginary components at different applied dc potentials at 3.2 Hz. As expected, the total cell admittance shows a peak at the equilibrium potential of the surface-bound redox probe. Removal of Rs from the impedance data results in the interfacial admittance, shown in Figure 5B. Nonfaradiac admittances (removal of the CPE) are then subtracted to isolate the real and imaginary faradaic admittances at various potentials, shown in Figure 5C. The faradaic admittance data of Figure 5C as fit to eq 4 to measure kct
Yfaradaic ) (Rct - j/ωCml)-1
(4)
Rct ) (RT/n2F2Akct)[R(Γoη-R + (1 - R)Γrη1-R]-1 (5) Cml-1 ) kctRct(η-R + η1-R)
(6)
Γo /Γr ) η ) exp[(nF/RT)(Eapp - E0)]
(7)
In eqs 4-7, kct is the charge-transfer rate constant, A is the electrode area, E0 is the formal potential, Γo and Γr are the surface concentrations of redox probes, and R is the transfer coefficient. All other variables have their standard meanings. Using Microsoft Excel’s goal seek function, a best fit of eq 4 to the data shown in Figure 5C resulted in the values kct ) 9 s-1, R ) 0.5, and Γtot ) 4 × 10-11 mol · cm-2. Note that the equivalent circuit used to derive eq 4 implies a simple series RC circuit and the experimental data, shown in Figure 3, requires an additional parallel resistor (Rd) to account for defects or ion
4920 J. Phys. Chem. C, Vol. 113, No. 12, 2009
Abhayawardhana and Sutherland
TABLE 1: Fitting Analysis Results of SAMs of 3 Using the Model Presented in Figure 4C at Selected Temperaturesa Cdl -2
Cml
temp (°C)
Rs (Ω · cm )
C (µF · cm )
n
Rct (Ω · cm )
C (µF · cm-2)
nb
Rd (kΩ · cm2)
χ2 c
kct (s-1)
7.3 16.1 24.8 34.6
2.1 3.1 1.3 2.4
21.8 (2) 27.3 (12) 34.9 (7) 30.4 (10)
0.85 0.86 0.84 0.88
696 (83) 382 (72) 98.4 (12) 58.1 (14)
34.0 (1) 18.0 (3) 60.9 (7) 30.5 (9)
0.90 0.92 0.91 0.93
326 (21) 84.7 (2) 61.0 (3) 22.3 (1)
0.67 0.29 0.74 0.50
1.1 1.9 7.4 12.6
2
b
2
a Numbers in parentheses are standard deviations from curve-fitting errors. b n is a power-law modifier in the CPE circuit element described by CPE ) 1/[|Z| · (jω)n]. c χ2 is a measure of the goodness of fit to the data using the equivalent circuit model shown in Figure 4c.
Figure 6. Bode plots of SAMs of 3 at 24.2 °C using frequencies ranging from 100 kHz to 10 mHz and applied dc potentials of -0.4, -0.14, and 0.3 V vs Ag|AgCl|KCl3M.
Figure 5. Admittance for SAMs of 3 at 3.2 Hz and 24.8 °C: (A) total cell admittance, (B) interfacial admittance, and (C) Faradaic admittance.
migration into the monolayer. Although there are differences between the equivalent circuit used by Brevnov et al. and that presented in Figure 4C, the charge-transfer rate determined using Brevnov et al.’s protocol is reasonable because the magnitude of Rd is large and will result in negligible contributions to the faradaic admittance. An Eyring plot of the kinetic data of SAMs of 3 from Table 1 is shown in Figure S4 of the Supporting Information. From the linear plot, Eact, ∆Hq, and ∆Sq of the charge-transfer reaction (kct) were found to be 69 kJ · mol-1, 67 kJ · mol-1, and -5 J · mol-1, respectively. To further understand the charge-transfer reaction, the electrodes were subjected to different dc applied potentials while the impedance spectrum was collected. The different applied potentials led to different impedance responses, which are
categorized into three regions: Eapp < E0, Eapp ) E0, and Eapp > E0. Representative Bode plots from each of those regions are shown in Figure 6. At anodic applied dc potentials (0.3 V), the phase angle of the Bode plots adopts a sigmoid shape that starts near 0° at high frequency and shifts close to 90° at frequencies of
