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Anion Effect on Mediated Electron Transfer through Ferrocene-Terminated Self-Assembled Monolayers G. Valincius,* G. Niaura,* B. Kazakevicˇiene˘ , Z. Talaikyte˘ , M. Kazˇeme˘ kaite˘ , E. Butkus, and V. Razumas Institute of Biochemistry, Mokslininku¸ 12, Vilnius 08662, Lithuania Received December 31, 2003. In Final Form: March 31, 2004 Inorganic anions strongly influence the electron transfer rate from the ascorbate to the ferroceneterminated self-assembled monolayer (SAM) composed of 9-mercaptononyl-5′-ferrocenylpentanoate (Fc(CH2)4COO(CH2)9SH, MNFcP). At the 1 M concentration level of the supporting anion (sodium salt electrolyte), a more than 10-fold increase in the electrocatalytic oxidation rate constant of the ascorbate is observed in the following sequence: PF6-, ClO4-, BF4-, NO3-, Cl-, SO42-, NH2SO3- (sulfamate), and F-. The sequence corresponds to the direction of increasing hydration energy of the corresponding anion, suggesting that highly hydrated ions promote electrocatalytic electron transfer to the ferrocene-terminated SAMs, while poorly hydrated ions inhibit it. Fourier transform surface-enhanced Raman spectroscopy (FT-SERS), in combination with cyclic voltammetry, indicates the formation of surface ion pairs between the ferricinium cation (Fc+) and low hydration energy anions, while, on the contrary, no ion pairs were observed in the electrolytes dominated by the high hydration energy anions. Though it is evident that the ion-pairing ability of hydrophobic anions is directly responsible for the electrocatalytic electron transfer inhibition, an estimate of the free, ion-unpaired Fc+ surface concentration shows that it cannot be directly related to the electron transfer rate. This suggests that the principal reason of the anion-induced electron transfer rate modulation might be related to the molecular level changes of the physical and chemical properties as well as the structure of the self-assembled monolayer.
Introduction Ferrocene-terminated self-assembled monolayers (FcSAMs) are one of the most studied redox-active twodimensional aggregates on metal surfaces. Starting from the pioneering work of Chidsey,1 these SAMs have been extensively used as convenient, robust, and well-reproducible surface self-assemblies for the kinetic and thermodynamic studies of electron transfer,2,3 the properties of the electrical double layer,4-6 and the influence of the redox microenvironment on electron transfer kinetics.5,7,8 One of the most noticeable features of the Fc-terminated SAMs is the sensitivity of their properties to the ionic composition of the solution.9,10 There is extensive experimental evidence10-14 indicating ion-pairing ability of the oxidized surface-bound ferricinium ion (Fc+) with anions, specifically perchlorate. It is also believed that ion pairing * Authors to whom correspondence should be addressed. Email:
[email protected] (G.V.);
[email protected] (G.N.). Phone: (3705)-272-9186. Fax: (3705)-2729196. (1) Chidsey, C. E. D. Science 1991, 251, 919. (2) Richardson, J. N.; Peck, S. R.; Curtin, L. S.; Tender, L. M.; Terrill, R. H.; Carter, M. T.; Murray, R. W.; Rowe, G. K.; Creager, S. E. J. Phys. Chem. 1995, 99, 766. (3) Weber, K.; Hockett, L.; Creager, S. E. J. Phys. Chem. B 1997, 101, 8286. (4) Creager, S. E.; Rowe, G. K. J. Electroanal. Chem. 1997, 420, 291. (5) Andreu, R.; Calvente, J. J.; Fawcett, W. R.; Molero, M. Langmuir 1997, 13, 5189. (6) Sondag-Huethorst, J. A. M.; Fokkink, L. G. J. Langmuir 1994, 10, 4380. (7) Creager, S. E.; Rowe, G. K. Anal. Chim. Acta 1991, 246, 233. (8) Creager, S. E.; Rowe, G. K. Langmuir 1993, 9, 2330. (9) Rowe, G. K.; Creager, S. E. J. Phys. Chem. 1994, 98, 5500. (10) Ju, H.; Leech, D. Phys. Chem. Chem. Phys. 1999, 1, 1549. (11) Ohtsuka, T.; Sato, Y.; Uosaki, K. Langmuir 1994, 10, 3658. (12) Kazakevicˇiene˘ , B.; Valincius, G.; Niaura, G.; Talaikyte˘ , Z.; Kazˇeme˘ kaite˘ , M.; Razumas, V. J. Phys. Chem. B 2003, 107, 6661. (13) Shimazu, K.; Ye, S.; Sato, Y.; Uosaki, K. J. Electroanal. Chem. 1994, 375, 409. (14) Sato, Y.; Mizutani, F.; Shimazu, K.; Ye, S.; Uosaki, K. J. Electroanal. Chem. 1999, 474, 94.
directly influences the stability of the Fc-terminated monolayer.13 For this reason, the vast majority of researchers employ perchlorate dominated buffers in their experiments.1-9,11-14 On the other hand, the absolute ionpairing constants are not known.15 For this reason, it is unclear how strong the surface ion pairing is. In other words, it is not known whether all Fc+ moieties, or just part of them, form tight ion pairs and how this factor might be related to the stability and other physical and chemical properties of Fc-terminated monolayers. Recently, we have reported an unusual ion-gating electrocatalytic effect induced by the ion pairing between perchlorate and surface-bound Fc+.12 In the case of the mediated electrocatalytic oxidation of the ascorbate, we observed an inhibition of the electrocatalysis in the solutions dominated by ClO4-. Conversely, the electrocatalysis was considerably facilitated as the perchlorate concentration was reduced,12 and the solution became dominated by the highly hydrated fluoride. The objective of the present work is to establish whether the observed modulation of the electrocatalysis rate is related only to a specific pair of anions, F- and ClO4-, or if it is more general and is due to the nature and the molecular level characteristics of the interactions between the redox SAM and the anion. Here, we investigate a series of buffered solutions containing the following anions: PF6-, ClO4-, BF4-, NO3-, Cl-, SO42-, NH2SO3- (sulfamate), and F-. Using the surface and molecular specific vibrational probe, Fourier transform surface-enhanced Raman spectroscopy16 (FT-SERS), and the cyclic voltammetry (CV) technique, we aimed to elucidate the molecular level picture of the electrocatalytical electron transfer phe(15) Rowe, G. K.; Creager, S. E. Langmuir 1991, 7, 2307. (16) (a) Tian, Z.-Q.; Ren, B.; Wu, D.-Y. J. Phys. Chem. B 2002, 106, 9463. (b) Niaura, G.; Gaigalas, A. K.; Vilker, V. L. J. Phys. Chem. B 1997, 101, 9250. (c) Meuse, C. W.; Niaura, G.; Lewis, M. L.; Plant, A. L. Langmuir 1998, 14, 1604. (d) Martusevicˇius, S.; Niaura, G.; Talaikyte˘ , Z.; Razumas, V. Vib. Spectrosc. 1996, 10, 271.
10.1021/la0364800 CCC: $27.50 © 2004 American Chemical Society Published on Web 07/10/2004
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nomenon. We believe that the possibility to modulate the electron transfer properties of the redox SAMs might be used in the development of electrocatalytic bioprocesses and electroanalysis as well as in the engineering of new SAM-based molecular devices.17-20 Experiments Chemicals. In all experiments, the Fc-SAMs were formed using 9-mercaptononyl-5′-ferrocenylpentanoate (Fc(CH2)4COO(CH2)9SH, MNFcP) synthesized as reported previously.21 Sodium salts were used throughout the work. Most salts, inorganic acids, alkali, and ascorbic acid were ACS reagent grade and were purchased from Sigma-Aldrich Chemie GmbH (Germany). Sulfamic acid and sodium tetrafluoroborate were of “chemically pure” category, and they were acquired from RIAP (Kiev, Ukraine) and Reachim (Moscow, Russia), respectively. The sodium hexafluorophosphate was from Strem Chemicals, Inc. (Newburyport, MA). Millipore purified (18.2 MΩ cm) water was used throughout the work. All experiments were carried out under argon (AGA, Sweden) flow. Electrode Preparation. For electrochemistry, the SAMs were prepared using polycrystalline gold electrodes (BAS, West Lafayette, IN) polished on alumina slurry (0.05 µm, Struers, Denmark). Right after the polishing, the electrodes were (i) sonicated for 10 min in a 1:1 v/v mixture of water and ethanol, (ii) etched in an aqua regia (a mixture of HCl, HNO3, and water in a 3:1:6 volume ratio) for 2 min and then sonicated for 10 min in water, and (iii) potentiostatically scanned (0.1 V/s) for ∼10 min in a 1.0 M sulfuric acid solution in the potential (E) window between -0.05 and 1.55 V (hereafter, the potentials are referred to the saturated sodium chloride calomel electrode (SSCE)). After the rinsing with water and ethanol, the electrodes were transferred in a 0.1 mM ethanol (95%) solution of MNFcP for 12-15 h. For the FT-SERS experiments, the electrodes were prepared in the same way as indicated above, except that the electrochemical roughening step was carried out in a 0.1 M KCl solution by the 50-fold E scanning (0.3 V/s) between -0.30 and 1.40 V, similarly as reported previously.12,22 All features of the voltammograms and electrocatalytic effects were qualitatively the same on the CV and FT-SERS electrodes. There were, however, some small quantitative differences observed for the differently prepared electrodes. Measurements. Electrochemical studies were carried out on a EG&G Versastat computerized potentiostat system (Princeton Applied Research, Princeton, NJ). The three-electrode conventional cell (10 cm3) was thermostated at 25 ( 0.2 °C. The platinum plate (∼2 cm2) served as an auxiliary electrode. As noted above, the SSCE was used as a reference electrode (E ) 239 mV vs standard hydrogen electrode). The real surface area of the working electrode was estimated from the integration of the Au surface oxidation charge (390 ( 10 µC cm-2).23 In the 1 M sulfuric acid solution, the current integration range spanned from ∼0.9 to 1.45 V. The Raman spectroscopic measurements were carried out with a FT-Raman spectrometer (PerkinElmer, Model Spectrum GX) equipped with an InGaAs detector operating at room temperature. An air-cooled diode-pumped Nd:YAG laser provided the excitation with an emission wavelength of 1064 nm. The laser beam was focused to a spot of an ∼1 mm2 area, and the laser power at the sample was set to 300 and 500 mW for the FT-SERS and solution FT-Raman measurements, respectively. To reduce photo- and thermoeffects, the spectroelectrochemical cell, together with the working electrode, was moved linearly with (17) Uosaki, K. Electrochemistry 1999, 67, 1105. (18) Campbell, D. J.; Herr, B. R.; Hulteen, J. C.; Van Duyne, R. P.; Mirkin, C. A. J. Am. Chem. Soc. 1996, 118, 10211. (19) Han, X.; Xu, G.; Dond, S.; Wang, E. Electroanalysis 2002, 14, 1185. (20) Alleman, K. S.; Weber, K.; Creager, S. E. J. Phys. Chem. 1996, 100, 17050. (21) Kazˇeme˘ kaite˘ , M.; Bulovas, A.; Smirnovas, V.; Niaura, G.; Butkus, E.; Razumas, V. Tetrahedron Lett. 2001, 42, 7691. (22) Gao, P.; Gosztola, D.; Leung, L.-W. H.; Weaver, M. J. J. Electroanal. Chem. 1987, 233, 211. (23) Burstein, R. K. Elektrokhimiya 1967, 3, 349.
Figure 1. Cyclic voltammograms of the MNFcP SAM on a gold electrode in the solutions containing different anions. The anion concentration is 0.1 M in a 0.01 M sodium phosphate buffer of pH 7.0. Potential scan rate, 100 mV/s; temperature, 25 °C respect to the laser beam (∼20 mm s-1).24 The experiments were carried out in a 180° geometry. The spectral resolution was set at 4 cm-1, and the wavenumber increment per data point was 1 cm-1. Typically, 700-1800 scans were coadded to enhance the signal-to-noise ratio of the FT-SERS spectra, while 200 scans were coadded for the solution FT-Raman spectra. The overlapped bands were deconvoluted into a sum of Gaussian and Lorentzian shapes by using the Grams/386 software package (version 2.04A, Galactic Industries Corp). None of the spectra presented were smoothed.
Results Figure 1 displays cyclic voltammograms (CVs) of the MNFcP SAM in 0.1 M solutions of different sodium salts at pH 7.0. As previously shown, 21 despite the presence of the ester group, MNFcP forms tightly (maximum surface coverage, ∼4.5 × 10-10 mol cm-2) packed monolayers that exhibit high stability in the perchlorate-based solutions. The stability of the MNFcP monolayers, however, was found to be much poorer in the solutions of Cl-, SO42-, F-, and NH2SO3- anions. In these solutions, the continuous potential cycling between -0.1 and +0.7 V induced an irreversible decrease of current peaks, which was shown to be related to the loss of the ferrocene headgroups in the nucleophilic attack by the solution anions.25 To minimize the experimental artifacts related to the monolayer instability, we limited the duration of the voltammetry experiments to two to four potential cycles, which is equivalent to 32-64 s of continuous polarization. (24) Niaura, G.; Gaigalas, A. K.; Vilker, V. L. J. Raman Spectrosc. 1997, 28, 1009. (25) (a) Popenoe, D. D.; Deinhammer, R. S.; Porter, M. D. Langmuir 1992, 8, 2521. (b) Zhang, L.; Godinez, L. A.; Lu, T.; Gokel, G. W.; Kaifer, A. E. Angew. Chem., Int. Ed. Engl. 1995, 34, 235. (c) Abbott, N. L.; Whitesides, G. M. Langmuir 1994, 10, 1493. (d) Prins, R.; Korswagen, A. R.; Kortbeek, A. G. T. G. J. Organomet. Chem. 1972, 39, 335.
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Table 1. Thermodynamic and Kinetic Data (Redox Conversion and Electrocatalytic Oxidation of Ascorbic Acid; Mean Values Are Shown (SE) of the MNFcP SAM in the Solutions of Various Anionsa anion
-∆Gs, kJ/mol
-
PF6 ClO4BF4NO3ClSO42NH2SO3F-
214b 254c 306b 347b 338c,d 472b
Em, mV vs SCE
fwhm, mV
Γ × 1010, mol/cm2
ks, s-1
kcat, M-1 s-1
252 ( 4 316 ( 5
92 ( 6 96 ( 8
4.45 ( 0.23 4.46 ( 0.20
112 ( 6 96 ( 6
318 ( 31 489 ( 59
357 ( 7 441 ( 6 457 ( 5 482 ( 13 499 ( 13 494 ( 13
89 ( 7 112 ( 8 129 ( 11 159 ( 10 179 ( 10 166 ( 17
4.43 ( 0.28 4.27 ( 0.24 2.71 ( 0.24 2.56 ( 0.37 1.25 ( 0.25 1.52 ( 0.34
88 ( 6 131 ( 5
748 ( 73 1336 ( 164 2325 ( 286 2464 ( 459 2728 ( 303 3931 ( 906
a Each solution contains 0.1 M sodium salt of the corresponding anion (1 M solutions were used for the determination of k ) and 0.01 cat M sodium phosphate buffer of pH 7.0. The ascorbic acid concentration was 0.001 M. b From ref 43. c From ref 44 d Value per 1 faraday.
In Figure 1, a clear trend is seen in the position of the anodic and cathodic peak potentials (Epa and Epc) as a function of the anion nature. In the sequence PF6- f ClO4f BF4- f NO3- f Cl- f SO42- f F- f NH2SO3-, the Epa and Epc values move anodically. In addition, the CV curves become noticeably asymmetric. As follows from columns 2 and 3 in Table 1, the position of the midpoint potential Em ) (Epa + Epc)/2 correlates fairly well with the sequence of the known hydration energies of the anions. Such a behavior was observed in other studies on different Fcterminated SAMs.9 Thus, it is logical to accept the position of Em as the measure of the anion “hydrophilicity/ hydrophobicity”, with PF6- being the most hydrophobic and F- and NH2SO3- being the most hydrophilic representatives. As the anion of a solution becomes more hydrophilic, one might clearly discern the additional redox peaks (marked by an arrow in Figure 1) gradually appearing on the CVs around 0.2 V. It seems as if, upon anodic shift of the main (anion-sensitive) redox peaks, the potential window around 0.20 V opens up for the scantily expressed redox peaks insensitive to the anion nature. These prepeaks are observed in all solutions which do not contain hydrophobic ions. They cannot be attributed to the phosphate species because they are present on the cyclic voltammetry curves even in the absence of the phosphates in the solution. The possible origin of these prepeaks was discussed earlier.12 One more trend is visible in Figure 1. The integrated voltammogram (current vs potential; j - E curve) area is almost constant in the cases of the most hydrophobic PF6-, ClO4-, BF4-, and NO3- anions, while it noticeably decreases in the solutions of hydrophilic anions (Cl-, SO42-, F-, and NH2SO3-). Conversely, the full width at halfmaximum (fwhm) increases (Table 1, column 4) in the same sequence indicating increased lateral repulsive interactions in the solutions of the hydrophilic anions. The electron transfer rate constant (ks) from the Fc headgroup to gold surface, as estimated using the Laviron method,26,27 was found to be approximately the same for the PF6-, ClO4-, BF4-, and NO3- ions with the values ranging from 88 to 131 s-1 (Table 1, column 6). On the other hand, at high potential scan rates (>1 V/s), we observed an ∼2 times smaller peak-to-peak potential separation, ∆Ep ) Epa - Epc, for the hydrophilic ions (data not shown). This suggests that the electron transfer between the Fc headgroup and the electrode occurs faster (26) Laviron, E. J. Electroanal. Chem. 1979, 101, 19. (27) As in the cases with most redox monolayers, we observed small (∼10 mV) initial peak-to-peak splitting, which does not disappear at the scan rates as low as 1-5 mV/s. The initial peak-to-peak separation was subtracted from the scan-rate-dependent peak-to-peak splitting when making estimates of the electron transfer rate constants for MNFcP monolayers.
Figure 2. Anodic scan of the background-subtracted voltammograms of ascorbic acid electrooxidation on the MNFcP modified gold electrode. The anion concentrations are 1.0 M except for fluoride; here, saturated NaF solution was used. All solutions were prepared in a 0.01 M sodium phosphate buffer of pH 7.0. The ascorbic acid concentration is 0.001 M. Potential scan rate, 100 mV/s; temperature, 25 °C.
in the presence of the hydrophilic anions. However, for these solutions, Laviron’s formalism is inapplicable because, as it is evident from Figure 1 and the data in columns 4 and 5 of Table 1, the monolayer isotherms do not exhibit Langmuirian properties. Therefore, for these ions, the values of the electron transfer rate constants are not presented in Table 1. Figure 2 shows the anodic scan of the backgroundsubtracted voltammograms of the ascorbic acid (AA) oxidation at the MNFcP SAM modified electrodes. These curves represent an electron transfer process occurring through the MNFcP monolayer via an electrocatalytic mechanism. There is an obvious dependence of the voltammogram shape and the current density magnitude on the anion. In the solutions of the most hydrophobic anions, PF6-, ClO4-, and BF4-, the voltammograms have a typical S-shaped character of the sluggish electrocatalytic process, while, starting with NO3-, the peak-shaped j - E curves appear, indicating the acceleration of the electrocatalytic electron transfer process. An important point is that the peak current (jp) values and the electrocatalytic oxidation rate constants of the AA (kcat) calculated in accordance with the Andrieux and Saveant method28,29 qualitatively follow the anion hydrophilicity (28) See the Supporting Information for details of the calculation of the electron transfer constant. (29) Andrieux, C. P.; Saveant, J. M. J. Electroanal. Chem. 1978, 93, 163.
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Figure 3. FT-SERS spectra of the MNFcP SAM on gold in the oxidized state at E ) 0.6 V (a) and the reduced state at E ) -0.1 V (b). The potentials are referred to the SSCE. A FT-Raman spectrum of a 0.5 M NaPF6 aqueous solution is also shown (c). The FT-SERS spectra were recorded in 0.01 M phosphate buffer containing 0.1 M NaPF6 (pH 7.0, 25 °C). For the FT-SERS and solution FT-Raman spectra, 700 and 200 scans were averaged, respectively. The excitation wavelength is 1064 nm. The laser power at the sample is 300 and 500 mW for the FT-SERS and FT-Raman experiments, respectively.
sequence accepted in the current work (compare columns 3 and 7 in Table 1). As evident from the kcat values, a more than 10-fold acceleration of the electrocatalytic electron transfer is observed upon the substitution of PF6- with F -. Figure 3 displays the FT-SERS spectra of the oxidized and reduced forms of the surface-bound MNFcP recorded in the PF6--containing solution at 0.6 and -0.1 V, respectively. From the spectra, it is seen that the bands of the reduced Fc ring at 318 and 1106 cm-1, which are assigned to the Fe-ring symmetric stretching νs(Fe-ring) and the ring symmetric stretching νs(ring) vibrations, respectively,12,30,31 transform to the bands at 311 and 1113 cm-1 upon the Fc headgroup oxidation (Figure 3, curve a). An additional clearly resolved peak is observed at 741 cm-1 in the FT-SERS spectrum of the oxidized MNFcP SAM (Figure 3, curve a). This band is absent in the solution without PF6-. Comparison of this spectrum with the FTRaman spectrum of a 0.5 M NaPF6 solution (Figure 3, curve c) immediately reveals that this spectral feature corresponds to the symmetric stretching vibration of PF6-,32 and indicates the uptake of the anion by the monolayer upon its oxidation, similarly as it has been observed in the solutions containing ClO4- anions.12,30 Figure 4 compares the difference FT-SERS spectra of the MNFcP SAM in the solutions containing various anions. These spectra were obtained as an algebraic difference between the spectrum recorded at E ) 0.6 V (30) Nishiyama, K.; Ueba, A.; Tanoue, S.; Koga, T.; Taniguchi, I. Chem. Lett. 2000, 930. (31) Nakamoto, K. Infrared and Raman spectra of inorganic and coordinated compounds; Wiley: New York, 1997. (32) Aroca, R.; Nazri, M.; Nazri, G. A.; Camargo, A. J.; Trsic, M. J. Solution Chem. 2000, 29, 1047.
Figure 4. Dependence of the difference (oxidized - reduced) FT-SERS spectra of the MNFcP SAM on the nature of the electrolyte anion. The spectra of oxidized and reduced states were recorded at 0.6 and -0.1 V (vs SSCE), respectively, in 0.01 M phosphate buffer (pH 7.0, 25 °C) containing 0.1 M NamX, where Xm- ) PF6-, ClO4-, BF4-, NO3-, SO42-, or NH2SO3-. The difference spectra were normalized to the oxidized MNFcP band at 1113 cm-1 (Figure 3a). The FT-Raman spectra of 0.5 M NamX aqueous solutions in the region of totally symmetric stretching vibration of anions are also shown (thin lines above the main spectra). For the FT-SERS and FT-Raman spectra, 700-1800 and 200 scans were averaged, respectively. The excitation wavelength is 1064 nm. The laser power at the sample is 300 and 500 mW for the FT-SERS and FT-Raman experiments, respectively.
(oxidized SAM) and the spectrum recorded at E ) -0.1 V (reduced SAM). Each FT-SERS spectrum of the MNFcP SAM is correlated with a small pattern of the FT-Raman spectrum recorded in a 0.5 M solution of the corresponding anion. These patterns are peculiar to the most intense totally symmetric stretching vibration of the selected anions.31,32 Thus, it is evident that some of the anions, namely, PF6-, ClO4-, BF4-, and NO3-, are capable of penetrating the MNFcP monolayer upon its oxidation, while the others, SO42- and NH2SO3-, are not. Moreover, Figure 4 demonstrates that the relative intensity of the SAM-bound anion peak does not match that recorded in the bulk solution. For example, the intensity of the waterdissolved NO3- vibrational band at 1048 cm-1 is signifi-
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Table 2. Relative Integrated FT-SERS Intensities of the Ion-Paired Anions (δX Values)a anion -
PF6 ClO4BF4-
δXb
anion
δXb
0.88 0.73 0.71
-
0.28 0 0
NO3 SO4NH2SO3-
a All the data were obtained in 0.1 M solutions of the corresponding sodium salts. b The SE values are not higher than 0.15.
cantly higher than that of PF6- at 743 cm-1. Nevertheless, the band intensities for the corresponding surface-bound anions are about the same, indicating that the same amount of the MNFc+P SAM (recall that the difference FT-SERS spectra are normalized to the intensity of the MNFc+P SAM peak at 1113 cm-1) binds a significantly lower amount of NO3- compared to PF6-. To quantify the ratio between the SAM-bound anion content and the surface concentration of the oxidized MNFcP SAM (δX), integrated spectral intensities of the values) were normalized to surface-bound anions (ASERS X the integrated spectral intensities of water-dissolved anions (AX values) and oxidized MNFcP (at 1113 cm-1; SERS AFc values). To take into account the difference in the + Raman cross section of anions, the spectra of 0.5 M NamX + 0.5 M NaClO4 solutions, where Xm- ) PF6-, BF4-, NO3-, SO42-, or NH2SO3-, were recorded, and the integrated intensity of the symmetric stretching vibrational band of each anion, Xm-, was normalized to the ClO4- peak intensity at 933 cm-1 (AX/AClO4). Thus, the δX parameter is determined by the following equation: SERS δX ) (ASERS /AFc + )(AClO /AX) X 4
(1)
The results of these calculations are summarized in Table 2. The variation of δX, which is proportional to the amount of ion-paired anions at the interface, with the nature of the anion reveals an interesting trend. In the sequence PF6- f ClO4- f BF4- f NO3-, the δX value progressively decreases, and, in the cases of SO42- and NH2SO3-, the δX parameter falls to about zero, as the bands of these two anions are totally absent in the FTSERS spectra (see Figure 4). Since the monolayer-bound (if present) F- and Cl- anions are “invisible” to the SERS technique, we cannot conclude whether these ions penetrate the SAM. However, by analogy, we might speculate that the ability of Cl- to ion pair should be somewhere between that of NO3- and that of SO42-, whereas highly hydrated F- should behave similarly to sulfate and sulfamate; that is, it should not ion pair with the Fc+. In addition to the above results, our FT-SERS studies revealed that, at relatively positive potentials, the oxidation of the MNFcP monolayer was not complete. Thus, Figure 5 shows the deconvolution of the FT-SERS spectrum band sensitive to the redox state of the Fc headgroup. As noted earlier, the two peaks at 1106 and 1113 cm-1 correspond to the symmetric stretching vibration of Fc rings in the reduced and oxidized states, respectively. The deconvolution of the band allows one to estimate the ratio between the integrated intensities of SERS SERS the oxidized (AFc values) and reduced (AFc values) + 0 forms of ferrocene at E ) 0.6 V. As can be seen from Figure SERS SERS depends on the nature of the 5, the ratio AFc + /AFc0 anions in solution. When an anion becomes more hydrophilic, the fraction of the oxidized ferrocene noticeably decreases. To transform the relative integrated FT-SERS intensities to the relative surface fraction of oxidized (θox) and reduced (θred) Fc moieties in the monolayer, the difference in the SERS cross section of the oxidized and
Figure 5. Dependence of the FT-SERS spectra of the MNFcP SAM on gold at E ) 0.6 V vs SSCE on the nature of the electrolyte anion. The spectra were recorded in 0.01 M phosphate buffer (pH 7.0, 25 °C) containing 0.1 M NamX, where Xm- ) PF6-, ClO4-, SO42-, or NH2SO3-. For the spectra, 800-1800 scans were averaged. The excitation wavelength is 1064 nm. The laser power at the sample is 300 mW.
reduced modes must be taken into account. Therefore, the correction factor γ was introduced: SERS* SERS SERS γ ) (AFc - AFc )/AFc 0 0 +
(2)
SERS* where AFc corresponds to the integrated intensity of 0 SERS the 1106 cm-1 band at E ) -0.1 V, while the AFc and 0 SERS AFc+ parameters represent the integrated intensities of the 1106 and 1113 cm-1 bands of the deconvoluted SERS* spectrum obtained at 0.6 V. Thus, the difference (AFc 0 SERS - AFc0 ) represents the integrated SERS intensity of the Fc groups that undergo complete oxidation at 0.6 V and SERS give the 1113 cm-1 band with integrated intensity AFc + . Now, the relative surface fraction of the oxidized MNFcP SAM can be estimated as follows:
SERS SERS SERS ] (3) θox/(θox + θred) ) (AFc + γ)/[(AFc+ γ) + AFc0
Figure 6 displays a relationship between θox/(θox + θred) and the nature of the dissolved dominant anion. A clear difference between two ion groups might be identified from these data. In the solutions of low hydration ions (PF6-, ClO4-, and BF4-), the surface Fc groups are almost completely oxidized at E ) 0.6 V. On the other hand, in the solutions of highly hydrated anions (SERS-“visible” NH2SO3- and SO42-), the fraction of the oxidized Fc headgroups falls below 0.4. Some intermediate situation is observed in the solution containing NO3- as a dominant anion. Discussion At pH 7.0, the electrocatalytic oxidation of AA at the gold electrode modified by the MNFcP SAM can be
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Figure 6. Dependence of the relative surface fraction of oxidized Fc groups (eq 3) in the MNFcP SAM on gold (E ) 0.6 V vs SSCE) on the nature of the electrolyte anion.
From this, we conclude that the formation of ion pairs inhibits electrocatalytic electron transfer, while the absence of ion pairing is a favorable factor for the process. It means that the necessary condition for effective electrocatalytic electron transfer through the Fc-terminated monolayer is the absence of the ion-pairing anions in a buffer. Only the four most hydrophobic anions, PF6-, ClO4-, BF4-, and NO3-, give clear spectroscopic evidence of their presence in the surface layer (see Figure 4). However, even within this group, the extent of the ion pairing differs. The data in Table 2 show that the relative intensities of the ion vibration bands (δX values) decrease ∼3-fold upon transition from the PF6- to NO3- ion. It means that only ∼1/3 of the oxidized MNFc+P surface cations are forming ion pairs with NO3- compared to PF6-. The electrocatalytic current density (Figure 2) and the corresponding electron transfer rate constant (kcat) (Table 1) increase to about the same extent. This allows hypothesizing that inhibition of the electrocatalytic electron transfer might be directly proportional to the ability of the anion to ion pair oxidized surface ferrocene moieties. In such a case, the rate of the AA electrocatalytic oxidation should be directly proportional to the surface concentration of free (non-ion-paired) surface MNFc+P ions. Consequently, the electrocatalytic current density could be written as
jcat ) nFkcatCAAΓFc+ presented by the following generalized 0
+
(6)
scheme:33
MNFc P SAM T MNFc P SAM + e
(4)
AA- + 2MNFc+P SAM f DAA + 2MNFc0P SAM + H+ (5) where AA- is the monoanionic form of AA (pKa ) 4.00)34 and DAA is the dehydroascorbic acid. From the results presented above, it follows that the nature of the solution-dominant anion considerably affects the electron transfer rate from AA to the surface-bound ferrocene groups of the MNFcP SAM. In the solutions of high hydration energy ions, Cl-, SO42-, NH2SO3-, and F-, the electrocatalytic oxidation of AA is much faster than that in the solutions of poorly solvated anions, PF6-, ClO4-, BF4-, and NO3-. What might be the reasons for such a conspicuous effect of ion nature on the electrocatalytic electron transfer rate? The FT-SERS spectra (see Figure 4) clearly affirm that only anions of low hydration energy penetrate the SAM upon oxidation of the ferrocene headgroups of the MNFcP SAM. According to many authors,9-11,35,36 such an uptake of ions by the Fc-terminated SAMs is related to the formation of tight (contact) ion pairs between the ferricinium (Fc+) and the anion. From our data, it is evident that not all anions are equally capable of forming contact ion pairs. The spectral bands of the hydrophilic representatives, for example, SO42- and NH2SO3-, which should exhibit a SERS signal, do not become visible upon oxidation of the monolayer (Figure 4). Consequently, they do not ion pair with the MNFc+P monolayer. At the same time, as it follows from Figure 2, these ions as well as Cl- and F- strongly promote the electrocatalytic oxidation of AA. (33) We claim that electron transfer occurs through the mediatory mechanism because, on the redox-inactive monolayers of similar length and structure, which have no ferrocene headgroups, we did not observe significant current flow upon addition of ascorbic acid to the solution. See the Supporting Information for the details. (34) Tur’yan, Y. I.; Kohen, R. J. Electroanal. Chem. 1995, 380, 273.
where n is the number of electrons, F is the Faraday constant, kcat is the rate constant of the electrocatalytic electron transfer reaction (eq 5), CAA is the bulk concentration of the ascorbic acid, and ΓFc+ is the surface concentration of the anion-unpaired ferricinium headgroups of the MNFcP SAM. As known from the literature,37-39 the rate of the oxidative electron transfer of AA might be controlled by either the first or second electron transfer step. Currently, we have no data which would allow us to determine the rate-limiting step in the systems under consideration. Therefore, here, the constant (kcat) is considered as an effective constant of the overall twoelectron transfer process represented by eq 5. From eq 6, it follows that the electrocatalytic current density should linearly depend on ΓFc+ if it is the only factor that determines the modulation of electrocatalysis by anions of different hydration energies. The correctness of this hypothesis could be assessed as follows. The data in Table 2 show that the relative surface concentration of NO3- is ∼3 times lower than that of PF6-. Although the absolute surface content of PF6- is not known, in the case of another hydrophobic representative, ClO4-, the quartz-microgravimetry data17,40 indicate that the surface concentration of the anion is close to the total amount of oxidized Fc. Moreover, our preliminary quartzmicrogravimetry experiments41 show that the same ionpairing stoichiometry is observed in the case of PF6-. This allows us to assume that, for PF6-, the fraction of anionunpaired MNFc+P is close to zero. Under this assumption, and taking into account the SERS data (see Table 2), in (35) Long, H. C. D.; Donohue, J. J.; Butry, D. A. Langmuir 1991, 7, 2126. (36) Kondo, T.; Okamura, M.; Uosaki, K. J. Organometal. Chem. 2001, 637-63, 841. (37) Ruiz, J. J.; Aldaz, A.; Dominguez, M. Can. J. Chem. 1977, 55, 2799. (38) Kulys, J.; Drungiliene, A. Electroanalysis 1991, 3, 209. (39) Hu, I.-F.; Kuwana, T. Anal. Chem. 1986, 58, 3235. (40) Shimazu, K.; Yagi, I.; Sato, Y.; Uosaki, K. J. Electroanal. Chem. 1992, 372, 117. (41) Data in preparation for publication.
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Figure 7. Schematic illustration of possible structural changes in the MNFcP monolayer upon transition from the solution of hydrophobic (left) to hydrophilic (right) anions.
the NO3--containing solution, the fraction of anionunpaired Fc+ headgroups should be ∼2/3 of the total amount of oxidized MNFcP. Because the spectroscopic experiments indicate no ion pairing, in the SO42- and NH2SO3dominated electrolytes, it follows then that compared to the nitrate-containing solution the concentration of the unpaired Fc+ headgroups in the solutions with SO42- and NH2SO3- should be higher by a factor of 1.5. To about the same factor the electrocatalytic current increases, as seen in the corresponding voltammograms of Figure 2. However, the above discussion did not take into account the fact that, in the solutions containing hydrophilic anions, the total amount of the oxidized MNFcP is significantly reduced (Figure 5). As is evident from Figure 6, at E ) 0.6 V, in the solutions containing SO42- or NH2SO3-, the fraction of the oxidized MNFcP form is at least 2-5 times lower than that in the solution containing NO3-. It obviously means that, in the case of SO42- or NH2SO3-, the absolute concentration of the anionunpaired MNFc+P is lower than that in the NO3dominated electrolyte. From that, it follows that the surface concentration of unpaired Fc+ headgroups does not influence directly the observed electrocatalytic effect. Consequently, the anions influence the electrocatalytic rate constant of AA oxidation through other factors. The electron transfer constant primarily depends on the particular transfer mechanism and reflects the molecular level phenomena occurring at the interface. As follows from our experimental data, there are a number of indications that the microenvironment of Fc+, the MNFcP SAM’s headgroup that directly participates in the electron transfer process, might be significantly different in the solutions with the hydrophilic and hydrophobic anions. First of all, in the solutions containing highly hydrophobic anions, the fwhm values of the voltammograms, which are close to 90 mV (see Table 1), indicate weak or no lateral repulsion between the surface MNFcP molecules in the oxidized state. In conjunction with the FT-SERS data, this fact suggests that the monolayer is electrically neutral. In other words, the interface between solution and the oxidized MNFcP SAM might be presented as a rigid two-dimensional ionic layer (Figure 7, left), in which strong (high-energy) ion pairing occurs. Consequently, high reorganization energy associated with the dissociation of the tight ion pair is expected.42 (42) Marcus, R. A. J. Phys. Chem. B 1998, 102, 10071.
On the other hand, solid-like monolayer structure would restrict thermal movement of the electron accepting Fc+ terminus, resulting in a relatively low transition state entropy of the electron transfer reaction activated complex. Hence, the electron transfer is expected to be much slower in the systems involving strong ion pairing. Conversely, the voltammetry data indicate strong repulsive interaction upon oxidation of the MNFcP SAM in the solutions of high hydration energy ions. As it is evident from Table 1, this interaction is reflected by the significantly broader fwhms and decreased integrated peak currents of the voltammograms (see the Γ values in Table 1). Low monolayer stability and high susceptibility of the ferrocene terminus to nucleophilic attack suggest that the monolayers are not electrically neutral. Given the fact that the FT-SERS data show no presence of the surface-bound hydrophilic anions (Figure 4), one might assume that the net positive charge of the monolayers is shielded by water dipoles and by diffusely distributed counterions, which are located outside the SAM. All these experimental facts lead to the conclusion that, in hydrophilic anion solutions, SAMs might be structurally less restrained than in the solutions with anions of high hydration energy (Figure 7, right). In the absence of tight ion pairs, the reorganization energy should be significantly lower, while a higher entropy could be expected for weakly structured monolayers. In our opinion, these factors cause significantly higher electrocatalytic electron transfer rate constants in the solutions of high hydration energy anions. Conclusions In this work, we present experimental evidence of the strong influence of anions on mediated electron transfer of ascorbic acid oxidation. It was found that, depending on the hydration energy of the anion in solution, the electrooxidation might be inhibited or accelerated. The hydrophilic (high hydration energy) anions promote the oxidation of ascorbate, while the hydrophobic (low hydration energy) anions strongly inhibit the electrocatalytic process. The ability to promote mediated electron transfer correlates with the values of the ionic hydration energy and increases in the following sequence of anions: PF6< ClO4- < BF4- < NO3- < Cl- < SO42- < NH2SO3(43) Marcus, Y. Ion Solvation; Wiley & Sons: New York, 1985. (44) Conway, B. E. Ion Hydration in Chemistry and Biophysics; Elsevier: New York, 1981.
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(sulfamate) < F-. Our data suggest that the surface concentration variation of the free (anion-unpaired) Fc+ headgroups does not directly influence the observed effects. However, it is obvious that an increase of the anion’s ability to ion pair with the ferricinium cation is followed by a decrease of the electrocatalytic current. Thus, the principal reason of the anion-induced electron transfer rate variation is the molecular level changes of the physical and chemical properties and, possibly, the structure of the self-assembled monolayer.
Valincius et al.
Acknowledgment. This work was partly supported by the Lithuanian Science and Studies Foundation under Contract No. T-24. Supporting Information Available: Details of the calculation of the mediated electron transfer rate constant, experimental material certifying the mediatory nature of the electron exchange between ascorbic acid and the electrode, and details on 9-mercaptononylhexanoate synthesis. LA0364800
