© Copyright 2000 American Chemical Society
SEPTEMBER 5, 2000 VOLUME 16, NUMBER 18
Letters Electrode Potential Effect on the Surface pKa of a Self-Assembled 15-Mercaptohexadecanoic Acid Monolayer on a Gold/Quartz Crystal Microbalance Electrode Kouki Sugihara and Katsuaki Shimazu* Division of Material Science, Graduate School of Environmental Earth Science, Hokkaido University, Sapporo 060-0810, Japan
Kohei Uosaki Division of Chemistry, Graduate School of Science, Hokkaido University, Sapporo 060-0810, Japan Received October 1, 1999. In Final Form: June 19, 2000 Surface mass titrations of a self-assembled monolayer (SAM) of 15-mercaptohexadecanoic acid on a gold/quartz crystal microbalance electrode have been performed at various electrode potentials. The apparent pKa of the monolayer, obtained from the pH value at half the total frequency change, shifted in a negative direction from 6.4 at -200 mV to 4.3 at -700 mV vs Ag/AgCl. When potentials more positive than 600 mV were applied, the pKa has a tendency to shift in the positive direction. These results demonstrate that the surface acid/base properties of the SAMs can be controlled by the electrode potential.
Introduction The evaluation of the acid/base properties of the selfassembled monolayers (SAMs) has recently received particular attention.1 The evaluation is essentially important because the surface properties of surface-immobilized molecules must be different from those in the bulk phase. Such a difference is due to the totally different environment around the molecule in the SAM from that in the bulk phase; for example, the electronic interaction between the molecule and the substrate will affect the electronic state of the acid/base center, and lateral attractive and/or repulsive interactions between the immobilized molecules will affect the dissociation of the * To whom correspondence should be addressed: KS e-mail:
[email protected]; phone: +81-11-706-2276; fax: +81-11-706-4868. (1) Finklea, H. O. In Electroanalytical Chemistry; Bard, A. J., Rubinstein, I., Eds.; Marcel Dekker: New York, 1990; Vol. 19, p 109.
surface acid. So far several different techniques have been applied to evaluate the acid/base properties including contact angle,2 capacitance,3 chemical or double layer force,4 spectroscopy,5,6 and the indirect laser-induced temperature jump method.7 The surface pKa of the (2) (a) Creager, S. E.; Clarke, J. Langmuir 1994, 10, 3675. (b) Lee, T. R.; Carey, R. I.; Biebuyck, H. A.; Whitesides, G. M. Langmuir 1994, 10, 741. (3) (a) Bryant, M. A.; Crooks, R. M. Langmuir 1993, 9, 385. (b) Wang, J.; Ward, M. D.; Ebersole, R. C.; Foss, R. P. Anal. Chem. 1993, 65, 2553. (c) Andreu, R.; Fawcett, W. R. J. Phys. Chem. 1994, 98, 12753. (4) (a) Vezenov, D. V.; Noy, A.; Rozsnyai, L. F.; Lieber, C. M. J. Am. Chem. Soc. 1997, 119, 2006. (b) van der Vegte, E. W.; Hadziioannou, G. Langmuir 1997, 13, 4357. (c) Hu, K.; Bard, A. J. Langmuir 1997, 13, 5114. (5) Cheng, S. S.; Scherson, D. A.; Sukenik, C. N. Langmuir 1995, 11, 1190. (6) (a) Sun, L.; Crooks, R. M.; Ricco, A. J. Langmuir 1993, 9, 1775. (b) Yu, H.-Z.; Xia, N.; Liu, Z.-F. Anal. Chem. 1999, 71, 1354. (7) Smalley, J. F.; Chalfant, K.; Feldberg, S. W.; Nahir, T. M.; Bowden, E. F. J. Phys. Chem. B 1999, 103, 1676.
10.1021/la991301t CCC: $19.00 © 2000 American Chemical Society Published on Web 08/04/2000
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mercaptoalkanoic acid SAM, which is most often examined, is dependent on the technique used; for example, for the mercaptoundecanoic acid SAM, the pKa ) 4.8-5.5 (adhesion force)4 and 6.7-11.5 (contact angle).2 In a previous paper,8a we have reported the evaluation of the SAMs of mercaptoalkanoic acids of different chain lengths by measuring the mass of the electrode surface during the titration of the SAMs using a quartz crystal microbalance (QCM). This technique was called “surface mass titration”. The resonance frequency decreased by ca. 2 Hz independent of the chain length. The mass increase is attributed to the cation association with the dissociated carboxylic acid. The surface pKa determined from the pH at half the total frequency change is 6.4-5.8 and slightly dependent on the chain length of the thiol; the longer chain mercaptoalkanoic acid gives a larger surface pKa. It should be mentioned that pioneering work using the QCM was conducted by Wang et al.9 They measured the frequency change of a 15-mercaptohexadecanoic acidcoated QCM electrode at various pH values to determine the surface pKa. Therefore, our idea is considered to be essentially the same as theirs. However, the results are quite different; they observed a large frequency increase of more than 1000 Hz during the titration and concluded that the frequency change is due to the viscoelastic change around the electrode. Unfortunately, we never observed such a large frequency increase although the surface mass titration experiments were repeated hundreds of times. Some differences in the detailed experimental conditions may lead to the different results. It would be more interesting if we could control the surface acid/base properties. This kind of study is particularly important for the design of functionalized surfaces. Only a few reports have described the electrode potential effect on the surface acid/base properties. White et al. have reported the reversible, electric-field-driven deprotonation of carboxylic acid groups in the mixed monolayer of 11-mercaptoundecanoic acid and 1-decanethiol on Ag(111), which gives rise to a current peak in the voltammogram.10 The electrode potential effect on the pKa of mercaptopyridine SAMs has been examined using FT-IR by Hara et al.;11 the apparent pKa shifted in a positive direction when negative potentials were applied. In this paper we describe the electrode potential dependence of the surface pKa of 15-mercaptohexadecanoic acid SAM on gold. To confirm the main results obtained by the surface mass titration method, the FT-IRRAS spectra were also taken.
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Figure 1. Surface mass titration curves for mercaptohexadecanoic acid SAMs at various electrode potentials: (a) -200 mV; (b) -300 mV; (c) -400 mV; (d) -500 mV; (e) -600 mV; (f) -700 mV. Solution was 0.1 M NaCl. character of Au(111) for the electrochemical formation/reduction of the surface oxide in 1 M H2SO4.13,14 The roughness factor of the gold was determined from the reductive charge of the surface oxide.13 A typical value is 1.3. The self-assembled monolayer of 15-mercaptohexadecanoic acid was constructed by immersing the gold substrate in a 1 mM ethanolic solution of the thiol for 12 h. Surface Mass Titration under Potential Control. The electrolyte solution was 0.1 M NaCl whose pH was initially adjusted to 4.0 with 0.1 M HCl. An aqueous solution of 0.1 M NaOH was added at a constant flow rate of 0.15 mL/min to the NaCl solution in which the QCM sensor head (MAXTEK model TPS550), a pH electrode, and electrodes for electrochemical measurements (reference electrode, Ag/AgCl (saturated NaCl); counter electrode, pt-Pt foil) were immersed. A personal computer was used to control the operation of the entire system. FT-IR Measurements. After the electrode potential of the SAM-coated gold electrode was held at various potentials for 10 min, the electrode was lifted out of the solution under the potential control and was quickly dried by flowing N2 gas. Infrared spectra were obtained in the external reflection mode using p-polarized light incident at 82° with a resolution of 4 cm-1 using a Biorad FTS 60A/816 spectrometer equipped with a liquid nitrogen cooled MCT detector. Typically, 1024 scans were averaged. Freshly prepared gold films were used to obtain the reference spectra.
Results and Discussion Experimental Section Materials. 15-Mercaptohexadecanoic acid (MHDA) was synthesized according to the procedure reported in the literature.12 NaCl (Wako), NaOH (Wako), and HCl (Kanto) were of reagent grade and used as received. All aqueous solutions were prepared with Milli-Q water and were sufficiently deaerated with 5N purity argon prior to use. Substrate and Monolayer Preparation. The substrate of the self-assembled monolayers was a 200 nm thin gold film, which was evaporated in a vacuum onto a 5 MHz, AT-cut quartz crystal wafer. The gold/quartz thus prepared showed the typical (8) (a) Shimazu, K.; Teranishi, T.; Sugihara, K.; Uosaki, K. Chem. Lett. 1998, 669. (b) Sugihara, K.; Teranishi, T.; Shimazu, K.; Uosaki, K. Electrochemistry 1999, 67, 1172. (9) Wang, J.; Frostman, L. M.; Ward, M. D. J. Phys. Chem. 1992, 96, 5224. (10) White, H. S.; Peterson, J. D.; Cui, Q.; Stevenson, K. J. J. Phys. Chem. B 1998, 102, 2930. (11) Hara, Y.; Wan, L.-J.; Taniguchi, I.; Osawa, M. ISE 49th Annual Meeting Extended Abstracts, 1998, p 201. (12) Troughton, E. B.; Bain, C. D.; Whitesides, G. M.; Nuzzo, R. G.; Allara, D. L.; Porter, M. D. Langmuir 1988, 4, 365.
Surface Mass Titrations at Constant Potentials. The titrations of a MHDA SAM were conducted at various electrode potentials. By plotting the frequency change, ∆f ()f(pH) - f(pH)4)) versus pH, we obtained the surface mass titration curves (Figure 1). The titration curve at -200 mV is essentially the same as that at open circuit previously reported,8 since the open circuit potential is close to this potential. The frequency decreased with the increase in the pH of the solution. This result does not agree with that previously reported by Wang et al.9 as described in the Introduction. They concluded from the impedance analysis of the quartz crystal that the shift is due to changes in the viscoelastic properties of the (13) Angerstein-Kozlowska, H. A.; Conway, B. E.; Hamelin, A.; Stoicoviu, L. J. Electroanal. Chem. 1987, 288, 429. (14) (a) Borges, G. L.; Kanazawa, K. K.; Gordon, J. G., II; Ashley, K.; Richer, J. J. Electroanal. Chem. 1994, 364, 281. (b) Watanabe, M.; Uchida, H.; Miura, M.; Ikeda, N. J. Electroanal. Chem. 1995, 384, 191. (c) Uosaki, K.; Ye, S.; Naohara, H.; Oda, Y.; Haba, T.; Kondo, T. J. Phys. Chem. 1997, 101, 7566.
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hydrodynamic layer in contact with the QCM. We also conducted an impedance analysis during the titration. The equivalent circuit resistance in which the change in viscoelastic properties is reflected15 does not change throughout the titration. Therefore, we concluded that the observed frequency decrease is due to the mass increase. The origin of this mass increase, we consider, is the ion association between the dissociated carboxylic acid or carboxylate anion and the sodium cation in solution, as described by eq 1 where the 1:1 stoichiometry is assumed.
Au-S(CH2)15COOH + NaOH f Au-S(CH2)15COO-Na+ + H2O (1) There exist several experimental results which support the ion association. First, any changes in frequency were not observed for the alkanethiol monolayers and a bare Au.8a Second, the total frequency change, ∆ft ()f(pH)8) - f(pH)4)) decreased or the total mass change increased linearly with the increase in the surface coverage of the 11-mercaptoundecanoic acid (MUA),8b which was estimated from the charge consumed for the electrochemical reductive desorption of the MUA. Third, both the magnitude of ∆ft and the pH region where the frequency decrease was observed are dependent on the cation used, strongly indicating that the cation is involved in the process.16 On the basis of the cation association, the total frequency change through the titration corresponds to the total number of carboxylates formed or carboxylic acids initially present. Assuming a 1:1 stoichiometry for the association between the cation and carboxylate and that sodium ion is not hydrated upon the association, the number of cations was calculated from the ∆ft value of 2.2 Hz to be 1.68 × 10-9 mol/cm2. This value is roughly within the range expected for a monolayer, but is higher than 7.72 × 10-10 mol/cm2 of the (x3×x3)R30° structure on Au(111).1,17 From the cation effects on the titration curve, we consider that the stoichiometric ratio (cation to carboxylate) is lower and that the sodium ion is hydrated. Although it is essentially important to determine the stoichiometry and the hydration number, this is beyond the scope of this paper and will be discussed in detail elsewhere.16 Since the frequency change is proportional to the associated cation or the number of carboxylate ions formed, the apparent pKa value is given by the pH at half the total frequency change. The apparent pKa at -200 mV is 6.4, which is the same as that observed at open circuit.8a The bulk-phase pKa for MHDA is not reported, but it seems to be close to the value of 4.6-5.0 for alkanoic acids with a long alkyl chain18 since the inductive effect by a thiol group is considered to be negligible for long-chain molecules. The surface pKa is, therefore, larger than that in the bulk phase. Such an alkaline shift in pKa has been obtained using other techniques: pKa values for the mixed SAMs of mercaptoundecanoic acid/alkanethiols by contact angle measurements are in the range of 6.2-11.5,2a and the pKa of mercaptopropionic acid by double-layer force measurements is 8.0.4c Some others have reported nearly identical pKa values with bulk phase values: pKa of (15) (a) Buttry, D. A.; Ward, M. D. Chem. Rev. 1992, 92, 1355. (b) Oyama, N.; Ohsaka, T. Prog. Polym. Sci. 1995, 20, 761. (16) Sugihara, K.; Shimazu, K.; Uosaki, K. In preparation. (17) (a) Widrig, C. A.; Alves, C. A.; Porter, M. D. J. Am. Chem. Soc. 1991, 113, 2807. (b) Kim, Y.-T.; Bard, A. J. Langmuir 1992, 8, 1096. (18) Serjeant, E. P., Dempsey, B. Eds. Ionization Constants of Organic Acids in Aqueous Solution; IUPAC Chemical Data Series; Pergamon Press: Oxford, New York, 1978.
Figure 2. Dependence of apparent pKa on the electrode potential.
mercaptoundecanoic acid SAM is determined from adhesion force measurements to be 4.8-5.5.4a,b The reason for the discrepancy in pKa is presently unclear. For the alkaline shift, several explanations have been proposed. We consider that the possible origin is the repulsive interaction between neighboring carboxylates because the apparent pKa of MUA SAM increased with the increase in the surface coverage of MUA in single-component SAMs and mixed SAMs of MUA with various alkanethiols.8b When the titration was conducted at potentials more negative than -200 mV, the frequency decrease was observed at the lower pH as shown in Figure 1b-f. The apparent surface pKa shifts in a negative direction with the decrease in the electrode potential and reaches 4.3 at -700 mV. The experiment at more negative potentials gives unreliable results due to the hydrogen evolution reaction. The total frequency change is 2.1 ( 0.1, independent of the electrode potential. We applied more positive potentials to the MHDA SAM-coated electrode during the titration. The apparent pKa remains unchanged up to 600 mV. At 700 mV, it starts to shift in a positive direction. However, titrations at potentials more positive than 800 mV are not allowed due to the oxidative desorption of MHDA from the electrode surface. The dependence of the apparent pKa on the electrode potential is summarized in Figure 2. These results indicate that the surface pKa is controllable in the range from 4.3 at -700 mV to 6.8 at 700 mV by the electrode potential. Frequency Response during the Potential Cycling. Instead of changing the solution pH by adding NaOH solution, the electrode potential was changed with a sweep rate of 5 mV/s under constant pH conditions. Figure 3 shows the frequency vs electrode potential curve at pH 5.4. On the negative scan, the frequency started to decrease at around -250 mV and reached a constant ∆f value of -2.1 Hz at -700 mV. On the reverse scan, the frequency increased and returned to the initial value at -100 mV. Therefore, the process is reversible (a hysteresis was observed when higher sweep rates were applied). The ∆f-E curve is easily understood based on the surface mass titration curves for the constant electrode potentials shown in Figure 1. At -200 mV, the frequency does not start to decrease at pH 5.4; in other words, the surface carboxylic acid does not dissociate. At more negative potentials, the frequency decreased because the apparent pKa shifts to lower values. At -700 mV, all the MHDA molecules in the SAM dissociate, and therefore, the frequency change reaches its maximum value, which corresponds to the total number of surface MHDA. For a more quantitative examination, ∆f ) f(pH)5.4) - f(pH)4) was calculated for all the titration curves in Figure 1 and is plotted in
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Figure 3. Frequency change in a MHDA SAM-coated quartz crystal during potential cycling at pH ) 5.4. Open circle shows ∆f ) f(pH)5.4) - f(pH)4) from Figure 1. Sweep rate was 5 mV‚s-1.
Figure 3 (open circles). These values almost lie on the ∆f-E curve. Therefore, both results, the surface mass titration curve at constant potential and the ∆f-E curve at constant pH, are in agreement with each other. Similar experiments were repeated but at different pH values. As expected, no frequency change was observed at pH 4.2 except at the negative potentials below -600 mV. At pH 6.4, half of the surface carboxylic acids dissociate at 0 mV because the pKa of the MHDA SAM is 6.4 at this potential (Figure 2). Therefore, the total frequency change in the ∆f-E curve at pH 6.4 is expected to be half of ∆ft (1.1 Hz). This was actually observed. At pH 7.5, most of the MHDA has dissociated even at -200 mV. Therefore, only a slight frequency change is observed. These potential dependencies of the frequency mean that the interfacial state is drastically changed by applying the electrode potentials (from uncharged to charged state or vice versa) particularly in the pH region close to the pKa. FT-IR Spectra at Various Electrode Potentials. To confirm the dissociation of the terminal carboxylic acid of SAM at negative potentials, the FT-IR spectra were taken ex situ after holding the electrode potential at various potentials for 10 min in 0.1 M NaCl (pH ) 5.3). The spectrum at 0 V shows a clear band at 1744 cm-1 characteristic of the carboxylic acid (CdO stretching of COOH)2a,5,12,19 (Figure 4a). A small and sharp band at 1440 cm-1 can be assigned to the CH2 scissors deformation.20 The bands attributed to the wagging modes that are expected for an all-trans conformation are obscure due to the relatively low signal-to-noise ratio in the wavenumber region for these modes (1400-1250 cm-1). When the electrode potential was shifted to -600 and -750 mV, a drastic change was observed in the spectra: the disappearance of the COOH stretching mode at 1744 cm-1 and the appearance of the COO- stretching modes at 1600 and 1420 cm-1 (asymmetric and symmetric stretching modes, respectively). The spectrum at -400 mV shows features of both COOH and COO-, although the COO- stretching modes are not very clear. These results indicate that the dissociation starts at least at -400 mV and is almost completed at -600 mV. This (19) (a) Smith, E. L.; Porter, M. D. J. Phys. Chem. 1993, 97, 8032. (b) Tao, Y.-T.; Hietpas, G. D.; Allara, D. L. J. Am. Chem. Soc. 1996, 118, 6724. (20) Porter, M. D.; Bright, T. B.; Allara, D. L.; Chidsey, C. E. D. J. Am. Chem. Soc. 1987, 109, 3559.
Figure 4. FT-IR spectra of MHDA SAM at various electrode potentials: (a) 0 mV; (b) -400 mV; (c) -600 mV; (d) -750 mV. The spectra were taken in dry air after holding the potential for 10 min.
potential region agrees with that where the frequency decrease was observed (Figure 3). Cheng et al. have obtained pH-difference attenuated total reflectance Fourier transform infrared (ATR-FTIR) spectra on open circuit for an alkyl carboxylate-bearing siloxane monolayer anchored to a germanium surface and determined the apparent pKa of the monolayer. Similar measurements and analyses under potential control are also possible. However, we reserve such a detailed analysis until the in situ IR measurements are conducted in order to compare the results from the IR spectra and the surface mass titration under the same conditions. Origin of pKa Shift by the Electrode Potential. As demonstrated using the QCM and Fourier transform infrared reflection adsorption spectroscopy (FT-IRRAS), the apparent pKa can be controlled by the electrode potential. It is very important to clarify the origin of this potential effect on the pKa. However, we have only a speculative idea on that at this moment. Before describing the idea, we will discuss the following two effects: local pH at the interface and the surface potential. Since we did not use a buffer solution to continuously
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change the solution pH, the local pH in the vicinity of the SAM may be different from the bulk pH. When negative potentials are applied, the local concentration of hydronium ion is expected to be higher. In these cases, the alkaline shift in the pKa should be observed. However, the results are reverse. The surface potential effect is recognized in several papers as the main factor of the positive shift in the pKa compared with that in the bulk phase. Smith and White first discussed this effect on the dissociation of the carboxylic acid SAM based on a simple dissociation model.21 Fawcett et al. modified their treatment by taking into account the partial compensation of the surface charge by the cation.22 According to their results, applying more negative potentials causes a negative shift in the surface potential. Since the apparent pKa, pKaapp, is given as a function of the surface potential, Ψ, by the equation
pKaapp ) pKa - (FΨ/2.3RT)
(2)
the negative shift in the surface potential leads to the positive shift in the apparent pKa. From the surface potential effect, therefore, a positive shift in the pKa is expected when negative potentials are applied. However, the present results do not agree with this prediction. Therefore, neither the local pH nor the surface potential becomes the main factor for the present pKa shift by the electrode potential. When negative potentials are applied, the local cation concentration as well as that of the hydronium ion is expected to be higher. This change in the local concentration of the cation causes the apparent pKa to shift in a negative direction. The true surface pKa is defined based on a simple dissociation of the surface carboxylic acid. Since the cation association simultaneously takes place with the dissociation, the apparent pKa we can determine is a function of the association constant, Kas, and the activity of the cation, aM+ (M+ presents the cation used) as expressed by the following equation assuming the 1:1 association between the cation and the surface carboxylate anion.23 (21) Smith, C. P.; White, H. S. Langmuir 1993, 9, 1. (22) (a) Fawcett, W. R.; Fedurco, M.; Kovacova, Z. Langmuir 1994, 10, 2403. (b) Andreu, R.; Fawcett, W. R. J. Phys. Chem. 1994, 98, 12753. (23) Jordan, C. E.; Corn, R. M. Anal. Chem. 1997, 69, 1449.
pKaapp ) pKa + pKas - log aM+
(3)
The dependence of the cation concentration on the apparent pKa was experimentally examined.16 With the increase in the concentration, the apparent pKa becomes lower as expected by eq 3. Therefore, the lower apparent pKa observed at more negative potentials may be attributed to the higher concentration of cation at the interface. A detailed examination of the cation and the cation concentration effects on the apparent pKa is now under way and will be reported elsewhere.16 In summary, the apparent pKa of the SAM of MHDA can be controlled in the range of 4.3 (-700 mV)-6.8 (700 mV) by changing the applied potential. There are few reports so far as to the potential dependence of the surface pKa. Hara et al. examined the surface pKa of 4-mercaptopyridine SAM on gold using the FT-IR method and found that the pKa shifts in a positive direction when negative potentials are applied.11 Such a shift can be explained by the surface potential effect. The voltammetric response observed for a Ag(111) coated with mixed SAMs of mercaptoundecanoic acid and decanethiol is also consistent with the prediction based on the electrostatic considerations by White et al.10 Contrary to these results, the direction of the pKa shift in this study is the reverse of that expected from the surface potential effect. As far as we know, this is first study to report the pKa shift of the SAM in a unique direction. This is of particular importance since it suggests that a factor other than the surface potential dominates the determination of the surface pKa, at least for this most popular acid SAM. It should be again emphasized that the interfacial state of the functional group is strongly dependent on the electrode potential; at pH ) 5.4, for example, the surface carboxylic acid is not dissociated at -200 mV, but is fully dissociated at -700 mV. This should be taken into account when we design and use the functionalized surface. Acknowledgment. This work is partially supported by a Grant-in-Aid for Scientific Research (09640711) and for Priority Area Research of “Electrochemistry of Ordered Interfaces” (09237101, 09237204, 10131203, 11118203) from the Ministry of Education, Science, Sports and Culture, Japan. LA991301T
