Indirect Laser-Induced Temperature Jump Study of the Chain-Length

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Langmuir 2003, 19, 9284-9289

Indirect Laser-Induced Temperature Jump Study of the Chain-Length Dependence of the pKa’s of ω-Mercaptoalkanoic Acid Monolayers Self-Assembled on Gold John F. Smalley* Materials Science Department, Brookhaven National Laboratory, Upton, New York 11973-5000 Received May 23, 2003. In Final Form: August 26, 2003 The indirect laser-induced temperature-jump (ILIT) technique is used to measure the pKa’s of selfassembled monolayers (SAMs on gold electrodes) composed of dibutanoic acid disulfide (DBAD) and 6-mercaptohexanoic acid (MHA) in contact with 1.00 M ionic strength NaClO4 electrolyte solutions. The ILIT technique may be used to determine the pKa of a surface-attached acid because the size of the ILIT response is related to the potential drop across the electrode/electrolyte interface that is, in turn, related to the extent of ionization of the acid. The pertinent data are the potentials of zero ILIT response (Epzr) versus pH. It is found that the measured Epzr versus pH data set for the DBAD monolayers is (within experimental error) the same as that measured for the MHA monolayers. Additionally, an analysis of the combined DBAD/MHA data set gives a pKa (4.5 ( 0.2) and a total concentration (ΓT) of either DBAD or MHA comprising the SAM (ΓT ) (5.4 ( 0.3) × 10-10 mol/cm2) that (again within experimental error) are the same as those previously determined for SAMs composed of 11-mercaptoundecanoic acid (MUA) in contact with 1.00 M ionic strength electrolyte solutions. This result implies that, at least in the range of thickness defined by monolayers made from DBAD and MUA, the properties of the microenvironment of the SAM/electrolyte interface are independent of monolayer thickness.

Introduction Self-assembled monolayers (SAMs) composed of thiols chemisorbed on various noble metals continue to be of both great scientific and technological interest because they effect highly ordered and stable surfaces.1,2 The chemical functionalities of these surfaces may also be easily controlled (i.e., engineered on the nanometer scale).3 Acidic (or basic) chemical functionalities are of particular importance in this regard because SAMs terminated by these functional groups provide a means of regulating the charge density at solid/liquid interfaces.4,5 The evaluation and comprehension of the properties of these interfacial charge densities are, in turn, relevant to a number of fields including adhesion,6 surface wetting,7 emulsion7 and colloid8 stability, biocompatibility,8 and the biophysical stability of membranes.9 The crystallization of inorganic salts,10,11 the preferential adsorption of metal ions,8 and the attachment and behavior of redox-active proteins such * E-mail: [email protected]. (1) Dubois, L. H.; Nuzzo, R. G. Annu. Rev. Phys. Chem. 1992, 43, 437. (2) Finklea, H. O. In Electroanalytical Chemistry; Bard, A. J., Rubinstein, I., Eds.; Marcel Dekker: New York, 1996; Vol. 19, pp 109335. (3) Giz, M. J.; Duong, B.; Tao, N. J. J. Electroanal. Chem. 1999, 465, 72. (4) Imabayashi, S.; Iida, M.; Hobara, D.; Feng, Z. Q.; Niki, K.; Kakiuchi, T. J. Electroanal. Chem. 1997, 428, 33. (5) Azzaroni, O.; Vela, M. E.; Martin, H.; Hernandez-Creus, A.; Andreasen, G.; Salvarezza, R. C. Langmuir 2001, 17, 6647. (6) Acid-Base Interactions: Relevance to Adhesion Science and Technology; Mittal, K. L., Anderson, H. R., Eds.; VSP: Zeist, The Netherlands, 1991. (7) Creager, S. E.; Clarke, J. Langmuir 1994, 10, 3675. (8) Schweiss, R.; Welzel, P. B.; Werner, C.; Knoll, W. Langmuir 2001, 17, 4304. (9) Scarlata, S. F.; Rosenberg, M. Biochemistry 1990, 29, 10233. (10) Li, J.; Lang, K. S.; Scoles, G.; Ulman, A. Langmuir 1995, 11, 4418. (11) Luo, J.; Kariuki, N.; Han, L.; Maye, M. M.; Moussa, L. W.; Kowalski, S. R.; Kirk, F. L.; Hepel, M.; Zhong, C. J. J. Phys. Chem. B 2002, 106, 9313.

as cytochrome c12-16 may also be studied using SAMs terminated by ionized carboxylic acids. The surface pKa of a SAM composed of a thiol terminated by an acidic functional group is one of the most important parameters controlling the surface properties of this monolayer.17-19 Measurements of these surface pKa’s, therefore, probe the microenvironments of the relevant SAM/aqueous electrolyte interfaces.8 The behaviors of the properties of these microenvironments as a function of the thickness of the SAM are also of great current interest.17,19 We have, for example, observed a surprising decrease in the activation energies (EA) of the interfacial electron-transfer reactions of redox couples irreversibly attached (tethered) to Au electrodes by both saturated20,21 and unsaturated22 tethers as a constituent of a SAM at short tether lengths (i.e., when the monolayer becomes thin). This decrease in EA could be due to a change in the microenvironment of the SAM/aqueous electrolyte interface as the thickness of the monolayer decreases. A number of techniques have been used to determine the pKa’s of monolayers composed of ω-mercaptoalkanoic (12) Song, S.; Clarke, R. A.; Bowden, E. F. J. Phys. Chem. 1993, 97, 6564. (13) Feng, Z. Q.; Imabayashi, S.; Kakuichi, T.; Niki, K. J. Electroanal. Chem. 1995, 394, 149. (14) Nahir, T. M.; Bowden, E. F. J. Electroanal. Chem. 1996, 410, 9. (15) Chen, X.; Ferrigno, R.; Yang, J.; Whitesides, G. M. Langmuir 2002, 18, 7009. (16) Leopold, M. C.; Black, J. A.; Bowden. E. F. Langmuir 2002, 18, 978. (17) Shimazu, K.; Teranishi, T.; Sugihara, K.; Uosaki, K. Chem. Lett. 1998, 669. (18) Kim, K.; Kwak, J. J. Electroanal. Chem. 2001, 512, 83. (19) Dai, Z.; Ju, H. Phys. Chem. Chem. Phys. 2001, 3, 3769. (20) Smalley, J. F.; Feldberg, S. W.; Chidsey, C. E. D.; Linford, M. R.; Newton, M. D.; Liu, Y.-P. J. Phys. Chem. 1995, 99, 13141. (21) Smalley, J. F.; Finklea, H. O.; Chidsey, C. E. D.; Linford, M. R.; Creager, S. E.; Ferraris, J. P.; Chalfant, K.; Zawodzinski, T.; Feldberg, S. W.; Newton, M. D. J. Am. Chem. Soc. 2003, 125, 2004. (22) Sikes, H. D.; Smalley, J. F.; Dudek, S. P.; Cook, A. R.; Newton, M. D.; Chidsey, C. E. D.; Feldberg, S. W. Science 2001, 291, 1519.

10.1021/la0348968 CCC: $25.00 © 2003 American Chemical Society Published on Web 10/02/2003

ω-Mercaptoalkanoic Acid Monolayers on Gold

acids on Au. These include voltammetric measurements,19,23,24 double-layer capacitance studies,25 contact angle titration measurements,7,8,26,27 piezoelectric quartz crystal microbalance measurements,17,28 streaming potential and current measurements,8 faradaic impedance measurements,18 atomic force microscope (AFM) measurements29,30 and the indirect laser-induced temperature jump (ILIT) technique31 employed in the present study. However, there continues to be considerable disagreement regarding both the absolute values of the pKa’s and, most especially, their behavior as a function of alkanethiol chain length measured by these various techniques.17,19 The ILIT technique has been used to determine the surface pKa’s (as a function of ionic strength) of a single chain length of an ω-mercaptoalkanoic acid (11-mercaptoundecanoic acid, MUA) monolayer in contact with NaClO4 (aqueous) electrolyte solutions.31 We found that the pKa decreased as the ionic strength of the electrolyte solution increased. This result was to be expected on the basis of diffuse double-layer effects (as described by GuoyChapman-Stern theory). Additionally, the total concentration of MUA comprising the monolayer determined by the ILIT technique was quite reasonable31,32 and was not found to be a function of ionic strength. In the present study, the ILIT technique is used to determine the pKa’s and carboxylic acid moiety surface concentrations (ΓT) of SAMs composed of 6-mercaptohexanoic acid (MHA) and dibutanoic acid disulfide (DBAD) at 1.0 M ionic strength. The alkanethiol chain-length dependence of the pKa’s of monolayers composed of ω-mercaptoalkanoic acids is, therefore, determined in the present study. It is found that the measured values of both pKa and ΓT are independent of chain length. The implications of this result in regard to the microenvironment of the carboxylic acidterminated SAM/aqueous electrolyte interface will be discussed. Experimental Section Equipment, Materials, and Methods. A detailed description of the ILIT technique, apparatus, cell, and experimental procedures has been given elsewhere.20,31,33-35 Briefly, in the ILIT technique, a small and rapid (10 ns) change in the open-circuit potential of an electrode/electrolyte interface is caused by a small (2-5 °C) indirect laser-induced temperature perturbation of this interface.33 The magnitude of the ILIT temperature change (∆T) is determined (using a piezoelectric transducer) from the size of the pressure pulse induced by the rapid interfacial temperature perturbation.20,35 Because the magnitude of the ILIT response is a function of the potential change across the SAM-modified electrode/electrolyte interface,31 the ILIT technique may be used (23) Bryant, M. A.; Crooks, R. M. Langmuir 1993, 9, 385. (24) (a) Smith, C. P.; White, H. S. Anal. Chem. 1992, 64, 2398. (b) Smith, C. P.; White, H. S. Langmuir 1993, 9, 1. (c) White, H. S.; Peterson, J. D.; Cui, Q.; Stevenson, K. J. J. Phys. Chem. B 1998, 102, 2930. (25) Kakiuchi, T.; Iida, M.; Imabayashi, S.; Niki, K. Langmuir 2000, 16, 5437. (26) Bain, C. D.; Whitesides, G. M. Langmuir 1989, 5, 1370. (27) Lee, T. R.; Carey, R. D.; Biebuyck, H. A.; Whitesides, G. M. Langmuir 1994, 10, 741. (28) Wang, J.; Frostman, L. M.; Ward, M. D. J. Phys. Chem. 1992, 96, 5224. (29) Hu, K.; Bard, A. J. Langmuir 1997, 13, 5114. (30) Kane, V.; Mulvaney, P. Langmuir 1998, 14, 3303. (31) Smalley, J. F.; Chalfant, K.; Feldberg, S. W.; Nahir, T. M.; Bowden, E. F. J. Phys. Chem. B 1999, 103, 1676. (32) Chidsey, C. E. D.; Loiacono, D. Langmuir 1990, 6, 682. (33) Smalley, J. F.; Krishnan, C. V.; Goldman, M.; Feldberg, S. W. J. Electroanal. Chem. 1988, 248, 255. (34) Smalley, J. F.; Geng, L.; Rogers, L.; Feldberg, S. W.; Leddy, J. J. Electroanal. Chem. 1993, 356, 181. (35) Feldberg, S. W.; Newton, M. D.; Smalley, J. F. In Electroanalytical Chemistry; Bard, A. J., Rubinstein, I., Eds.; Marcel Dekker: New York, 2003; Vol. 22, pp 101-180.

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Figure 1. Examples of ILIT responses obtained from Au electrodes coated with either DBAD or MHA in contact with 1.00 M NaClO4 electrolyte solutions at various pH’s. (O) DBAD monolayer, pH ) 2.28, Ei ) -100 mV vs SSCE, ∆Teq ) 4.2 K; the solid line describes the fit of these data to eq 7 for a purely thermal response (B ) 0) where A ) -0.55 mV. (3) DBAD monolayer, pH ) 4.40, Ei ) +300 mV vs SSCE, ∆Teq ) 4.7 K; the solid line describes a purely thermal response fit to eq 7 (B ) 0) where A ) 0.54 mV. (4) MHA monolayer, pH ) 7.10, Ei ) +100 mV vs SSCE, ∆Teq ) 4.3 K; the solid line describes a purely thermal response fit to eq 7 (B ) 0) where A ) -2.66 mV. (]) MHA monolayer, pH ) 6.68, Ei ) +500 mV vs SSCE, ∆Teq ) 4.1 K; the solid line describes a fit of these data to eq 7, where km ) 2.0 × 107s-1, A ) -0.65 mV, and B ) 0.63 mV. to determine the pKa of a surface-confined acidic functional group; see the description of the data analysis given below. As before,20-22,31,34 the gold-film electrodes used in the present study were vapor deposited over a titanium underlayer (∼500-Å thick) on quartz substrates (1-in.-diameter disks made of Homosil quartz by Heraeus Amersil). The total (gold plus titanium) thickness of these electrodes was ∼1.0 µm. The titanium underlayer effects a uniform orientation of the 111 facets of the gold microcrystallites comprising the vapor-deposited gold layer.32 The area of these electrodes in contact with the electrolyte solutions was 0.71 cm2, and they were cleaned in an argon plasma before use. The dibutanoic acid disulfide36 was a gift from Professor Edmond F. Bowden (North Carolina State University), and the 6-mercaptohexanoic acid was a gift from Professor Harry O. Finklea (West Virginia University). A clean electrode was placed in an ethanol solution containing either 1.0 × 10-3 M dibutanoic acid disulfide or 1.0 × 10-3 M 6-mercaptohexanoic acid for approximately 3 days. (As before,31 no alkanethiol diluent was used in these experiments.20,26) The electrode was taken out of this solution, rinsed in neat ethanol, quickly dried in a stream of Ar, and attached to the ILIT cell31 that was to contain the electrolyte solution. Initially in each ILIT experiment, this aqueous electrolyte solution contained 1.00 M NaClO4 along with ∼0.01 M H3PO4 that produced a pH of ∼2.00. Cyclic voltammograms of the acidterminated SAMs taken at this point in the ILIT experiments were similar to those shown in Figure 1a of ref 31. (The doublelayer charging currents observed in these cyclic voltammograms are somewhat (less than a factor of 2) larger than those measured previously (for 11-mercaptoundecanoic acid monolayers in contact with 1.00 M ionic strength electrolyte solutions).31 However, the total double-layer capacitances (CT) determined from these (36) The formation and exchange kinetics of the adsorbed species37,38 as well as the position of the X-ray photoelectron spectroscopy 2p peak of the adsorbed sulfur37 demonstrate that both thiols and disulfides react with gold surfaces to form the same species.

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currents are still very much less than the diffuse double-layer components of these CT (see the Data Analysis section below).) During the ILIT experiments, the pH of the electrolyte solutions was raised by the addition of a 1.00 M NaOH solution and was lowered by the addition of a 1.00 M HClO4 solution. The pH of these electrolyte solutions was determined during the course of the ILIT experiments by removing a 0.5-mL aliquot of the solution from the ILIT cell and measuring its pH with a pH meter.31 Johnson Matthey 99% pure (metals basis) NaClO4‚H2O, Baker reagent-grade phosphoric acid, Mallinckrodt reagent-grade 2.0 N sodium hydroxide, and Baker Ultrex perchloric acid were all used as received to make the solutions used in the present study. Water was purified in a Millipore Mill-Q Plus system. All ILIT experiments were performed at room temperature (23 ( 1 °C), and a saturated sodium calomel reference electrode (SSCE) was used in all experiments. Data Analysis. As before,31 the SAM-modified electrode/ electrolyte interface is assumed to be adequately described by the two-layer model of Smith and White.24b In this model, all acidic moieties are assumed to lie in a single plane of acid dissociation (PAD) at the terminus of the SAM, and the diffuse double layer starts at the PAD and extends into the electrolyte solution. The first layer, therefore, consists of the entire SAM up to (and including) the ionizable functional group, and the second layer is the diffuse double layer. The open-circuit potential (E) between the gold electrode and the (SSCE) reference electrode is then given by

σΜ σA/B E) + + VD CT Cdl

(1)

where CT (F/cm2) is a series combination of the integral capacitance associated with the monolayer up to the PAD (Cfilm, F/cm2) and the integral capacitance associated with the diffuse double layer (Cdl, F/cm2), σM is the surface concentration of charge on the gold electrode, σA/B is the concentration of charge associated with the fraction of the acid that is ionized, and VD is the dipole potential associated with the entire system including any dipoles within the monolayer. From eq 1, the response to the ILIT temperature perturbation (∆E/∆T) is determined to be31

d ln[CT] ∆E ) -(Ei - VD) ∆T dT σA/B d ln[Cdl] d ln[CT] d lnKD [H+] + Cdl dT dT dT [H ] + KD

(

(

))

dVD (2) + dT

where Ei is the initial potential set with the ILIT apparatus potentiostat,34 the quantity dVD/dT includes the Soret potential induced by the temperature difference between the electrode and the bulk (electrolyte) solution,39 and KD is the dissociation constant of the surface-attached acid. Because CT/Cdl , 1.0, KD is given by31

KD ) K∞ exp

[ ] FσA/B RTCdl

(3)

where K∞ would be the dissociation constant of the acid at infinite ionic strength (where Cdl would also be infinite40). Note that, according to eq 3, KD has no explicit dependence upon the potential (E) and σM so that

[(

pH ) pK∞ - log10 -

)]

ΓTF +1 σA/B

-

FσA/B 2.3RTCdl

(4)

(37) Bain, C. D.; Troughton, E. B.; Tao, Y.-T.; Evall, J.; Whitesides, G. M.; Nuzzo, R. G. J. Am. Chem. Soc. 1989, 111, 321. (38) (a) Biebuyck, H. A.; Whitesides, G. M. Langmuir 1993, 9, 1766. (b) Biebuyck, H. A.; Bain, C. D.; Whitesides, G. M. Langmuir 1994, 10, 1825. (39) Smalley, J. F.; MacFarquhar, R. A.; Feldberg, S. W. J. Electroanal. Chem. 1988, 256, 21. (40) Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Fundamentals and Applications; Wiley & Sons: New York, 1980; p 507.

where ΓΤ (mol/cm2) is the total surface concentration of (both ionized and neutral) carboxylic acid moieties and σA/B is negative in the present study. The pKa of the surface-attached acid is defined as the value of the pH of the electrolyte solution in contact with the SAM when half of the acid functional groups is ionized.27 The potential (Epzr) at which the ILIT response (∆E/∆T) is zero is independent of the determination of ∆T described above. The most accurate method of measuring the pKa of a surfaceattached acid, therefore, would be to determine Epzr as a function of pH (i.e., from eqs 2 and 3).

Epzr ) VD +

[

dVD/dT d ln[CT]/dT d ln[Cdl]/dT

-

(

σA/B d ln KD/dT [H+] -1Cdl d ln[CT]/dT d ln[CT]/dT [H+] + KD

)]

(5)

where

d lnKD ) dT

( )( [ ( )(

)

FσA/B 1 d ln[Cdl] d lnK∞ + dT RTCdl T dT +

)]

FσA/B [H ] 1RTCdl [H+] + KD

(6)

and because σA/B is a function of pH, Epzr will vary with pH. If σA/B (and, consequently, KD), d ln K∞/dT, d ln[CT]/dT, d ln[Cdl]/dT, Cdl, VD, and dVD/dT are all independent of potential, then a plot of ∆E/∆T versus Ei (eq 2) will be linear. Epzr may be conveniently obtained from such a plot (as its intercept with the potential axis) even when Epzr is not in an accessible range of potential.31 The values of Epzr so obtained were then fitted to a combination of eqs 4-6, resulting in values of VD, dVD/dT, ΓT, K∞, d lnK∞/dT, and pKa. For these fits, Cdl and d ln[Cdl]/dT were calculated using Gouy-Chapman theory modified for the onset of dielectric saturation41 caused by the large concentration of charge (σA/B) that exists at the SAM/electrolyte solution interface when a significant fraction of the carboxylic acid moieties are ionized. (See Figures 2 and 3 in ref 31.) The ILIT responses (∆V(t)) observed in the present study were fitted to20,34

∆V(t) ) A∆T*(t) + kmB

∫e

t -k (t - τ) m

0

∆T*(τ) dτ

(7)

where A∆T*(t) would be the (open-circuit) response at time t in the absence of any relaxation, B∆T*(t) would be the amplitude of the relaxation at time t if km were infinite, and km is the measured first-order rate constant for the relaxation. ∆T*(t) is defined as

∆T*(t) )

∆T(t) ∆Teq

(8)

where ∆T(t) is the convolution of the interfacial temperature change at time t and the instrument response function, and ∆Teq is the interfacial temperature change that would be produced if all of the absorbed heat were uniformly distributed in the electrode and none of this heat was lost to either the dielectric “backing” material or the electrolyte solution. The total response (defined as ∆E/∆Teq) of an electrode/electrolyte interface to an ILIT temperature perturbation is the sum of the quantities A/∆Teq and B/∆Teq so that ∆E (in eq 2) is (A + B) and31

(A + B) ∆E ) ∆Teq ∆Teq

(9)

Results and Discussion The cyclic voltammograms described in the Experimental Section demonstrate that the SAMs investigated in the present study effect an almost completely (see below) (41) Grahame, D. C. J. Chem. Phys. 1950, 18, 903.

ω-Mercaptoalkanoic Acid Monolayers on Gold

Figure 2. Examples of plots of ∆E/∆Teq (see the text) vs Ei. (O) DBAD monolayer, pH ) 1.98. (b) MHA monolayer, pH ) 3.01. (3) DBAD monolayer, pH ) 4.40. (1) MHA monolayer, pH ) 6.01. (×) MHA monolayer, pH ) 7.31. (4) DBAD monolayer, pH ) 8.66.

insulating layer of low dielectric constant material between the gold electrodes and the electrolyte solutions. The charging currents observed in these cyclic voltammograms are somewhat larger because the thinner monolayers investigated here are slightly more permeable to ions,2,31,32,42,43 and the simple fact that the dielectric layers are thinner would also result in larger charging currents.2,20 Examples of the ILIT responses observed in the present study are shown in Figure 1. Most of these responses are well fit by eq 7 with B set equal to 0 (i.e., the time dependence of the ILIT response follows that of the thermal perturbation). However, when Ei g 500 mV versus SSCE (for pH < 6.0) or when Ei > 200 mV versus SSCE (for pH > 6.0), the measured ILIT responses do contain a relaxation. (See Figure 1.) As observed with the MUA monolayers,31 the size of this relaxation is usually quite small, but for pH > 6.0, it increases with increasing potential. The measured rate constant for this relaxation is quite similar to that for MUA monolayers (i.e., km ) (1.2 ( 0.2) × 107s-1 at 1.0 M ionic strength),31 being (1.1 ( 0.3) × 107 s-1 for the monolayers made from DBAD and (1.9 ( 0.4) × 107 s-1 for the monolayers made from MHA. As was also observed with the MUA monolayers, the relaxation rate constants measured in the present study are not functions of potential. These observations are consistent with the conclusion31 that the relaxations of the ILIT responses observed in the present study are (also31) due to a comparatively slow relaxation of the monolayers’ film capacitance (Cfilm) and with the additional observation that monolayers composed of ω-mercaptoalkanoic acids are quasi-liquid.32 Figure 2 contains examples of the plots of ∆E/∆T versus Ei generated in the present study. The linear ranges in all of these plots are at least from -100 to +200 mV versus SSCE.31 For these linear ranges of potential, the slopes of the fitted lines are equal to the quantity -d ln[CT]/dT (eq 2). Figure 3 contains plots of d ln[CT]/dT versus pH for monolayers made from both DBAD and MHA. Within experimental error, the data plotted in Figure 3 for the DBAD monolayers may be superimposed on that for the MHA monolayers. The behavior of the combined DBAD/

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Figure 3. Plots of d ln[CT]/dT (eq 2) vs pH for both (O) DBAD and (b) MHA monolayers. The solid-line curve describes a fifth power polynomial fit to the combined DBAD/MHA data set. This curve has no theoretical significance. The dotted-line curve describes the fourth power polynomial fit to the d ln[CT]/dT vs pH data obtained from Au electrodes coated with MUA, which were in contact with 1.00 M ionic strength electrolyte solutions.31

Figure 4. Plots of Epzr (see eq 5) vs pH for both (O) DBAD and (b) MHA monolayers. The solid-line curve describes the fit of the combined DBAD/MHA data set to eqs 4-6 (Table 1). The dotted-line curve describes the fit of the entire (including data for pH > 7.5) Epzr vs pH data set obtained from Au electrodes coated with MUA, which were in contact with 1.00 M ionic strength electrolyte solutions.31

MHA data set shown in Figure 3 is somewhat different from that observed with the MUA monolayers investigated previously.31 (See the dotted curve in Figure 3.) However, as has been explained previously,31 a complete interpretation of this behavior would require detailed information concerning the structure of all of these monolayers as a function of temperature and pH. The ILIT technique does not provide such information, so such an interpretation is beyond the scope of the present study.31 Plots of Epzr versus pH are shown in Figure 4 for monolayers made from both DBAD and MHA. The behavior of the data shown in Figure 4 (as well as the data shown in Figure 3 and the relaxation rate constant) is

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Table 1. Results from the Fit of the Epzr versus pH Dataa,b parameter

DBAD/MHAc

MUAd

pKa pK∞ ΓT/(×10-10mol/cm2) VD/mV vs SSCE (dVD/dT)/(mV/K) (d ln K∞/dT)/10-2 K-1

4.5 ( 0.2 3.0 ( 0.2 5.4 ( 0.3 74 ( 67 -0.02 ( 0.06 -0.4 ( 0.3

4.4 ( 0.2 2.9 ( 0.2 5.2 ( 0.3 83 ( 65 0.07 ( 0.04 0.1 ( 0.6

a See Figure 4. b Data obtained at 1.0 M ionic strength. c Results from the fit (to eqs 4-6) of the combined DBAD/MHA data set. d From the fit of the entire 1.0 M ionic strength MUA data set.31

reproducible from experiment to experiment (monolayer to monolayer) and is independent of the pH and potential history of a particular monolayer. Also note that, as with the d ln[CT]/dT versus pH data shown in Figure 3, the Epzr versus pH data sets obtained for each of the two monolayer thicknesses investigated in the present study may (within experimental error) be superimposed upon each other. The combined (DBAD/MHA) data set shown in Figure 4 was fitted to eqs 4-6. (See the solid curve in this Figure.) The values of the parameters (pKa, pK∞, ΓT, VD, dVD/dT, and d lnK∞/dT) obtained from this fit as well as the values of the same parameters obtained previously for MUA monolayers (at 1.0 M ionic strength)31 are given in Table 1. (The error bounds reported in Table 1 come from a statistical analysis of the fit described in Figure 4.) The most important thing to note about the data reported in Table 1 is that the pKa (and, consequently, the pK∞) and the total surface concentration of carboxylic acid moieties for SAMs made from both DBAD and MHA are the same as those measured for SAMs made from MUA. Additionally, the quantities VD, dVD/dT, and d lnK∞/dT reported in Table 1 for the DBAD and MHA monolayers are the same (within experimental error) as those determined for MUA monolayers. However, the data plotted in Figure 4 provide no indication of the Ca2+ cation complex formation44 that is evident in the MUA monolayer Epzr versus pH data when pH > 7.5. (See the dotted curve in Figure 4 and Figure 11 in ref 31.) Apparently then, the equilibrium constant for this complex formation is a function of the thickness of the monolayer (i.e., it decreases as this thickness decreases). The difference between the dottedline curve (for MUA SAMs) and the solid-line curve (for DBAD/MHA SAMs) in Figure 4 is due to both this lack of Ca2+ cation complex formation at high pH’s on the DBAD/MHA monolayers and the dissimilar behavior of d ln[CT]/dT (as a function of pH) for the DBAD/MHA monolayers compared to that for MUA monolayers. (See Figure 3.) It should be emphasized that the (ILIT) measured ΓT reported here and in ref 31 are not only independent of the thickness of the monolayer but also quite reasonable. The experiments of Chidsey and Loiacono32 as well as the cyclic voltammetric results reported here and in ref 31 indicate that monolayers composed of ω-mercaptoalkanoic acids are not densely packed. The measured ΓT, therefore, should be somewhat less than that expected for a densely packed (x3 × x3)R30° structured alkanethiol monolayer on Au(111) (i.e., 7.7 × 10-10mol/cm2),2,45,46 which is exactly what is observed. Reductive desorption4 and quartz crystal (42) Porter, M. D.; Bright, T. B.; Allara, D. L.; Chidsey, C. E. D. J. Am. Chem. Soc. 1987, 109, 3559. (43) Rebek, J., Jr.; Duff, R. J.; Gordon, W. E.; Parris, K. J. Am. Chem. Soc. 1986, 108, 6068. (44) The same batch of NaClO4‚H2O was used in the experiments performed in the present study and those reported in ref 31 so that the 1.0 M electrolyte solutions used in all of these experiments contained the same (small) number of Ca2+ cations as an impurity.

microbalance (QCM)17 experiments indicate that the ΓT for ω-mercaptoalkanoic acid SAMs is 2 to 3 times that measured in the present study and the study described in ref 31. However, systematic errors2 present in both the reductive desorption and QCM techniques (e.g., uncertainties associated with double-layer charging current subtraction in the reductive desorption technique2) may very well account for these differences in the measured values of ΓT. Additionally, evaluations of ΓT using the AFM technique29,30 are considerably (as much as a factor of 100)30 smaller than those values reported here and in ref 31. There are a number of possible explanations for the lowerthan-expected (and physically reasonable) ΓT measured in these AFM experiments,30,47 but the invariance of the ILIT-measured ΓT as a function of ionic strength31 demonstrates that the explanation cannot be a failure of Gouy-Chapman-Stern theory to describe the charge distribution adequately in the double-layer48 or counterion (Na+) adsorption that screens the surface charge.30,47 Despite the inconsistency between the value of ΓT measured by Hu and Bard29 and the value of ΓT measured in the present and previous31 ILIT studies, the value of pKa reported by Hu and Bard29 is completely consistent with that reported here and in ref 31. That is, in ref 29, the pKa reported for 3-mercaptopropionic acid monolayers in contact with 1.0 × 10-3 M ionic strength electrolyte solutions is 7.7. A SAM pKa decreases by 1 unit for every factor of 10 increase in the ionic strength of the electrolyte solution in contact with the SAM.31 The pKa of a 3-mercaptopropionic acid SAM should then be 4.7 at 1.0 M ionic strength of the electrolyte solutionsthe same, within experimental error, as that given in Table 1 for DBAD, MHA, and MUA monolayers. Additionally, Shimazu et al. have reported that the pKa’s of ω-mercaptoalkanoic acid SAMs on gold (at 0.1 M ionic strength) are only slightly dependent upon the length of the alkane chain, but these pKa’s, when they are converted to their values at 1.0 M ionic strength, are as much as 1.0 unit larger than that reported here and in ref 31. One possible explanation for the lack of consistency in the pKa’s measured in various laboratories is the variety of gold substrate preparations used in these laboratories. However, for the argon ion plasma-cleaned (approximately) 111 gold surfaces used in our laboratory,34 the lack of variation (again within experimental error) in the values of VD, dVD/dT, and km (the ILIT transient relaxation rate constant) indicates that the dielectric properties of ω-mercaptoalkanoic acid SAMs are not dependent upon the chain length of the thiol. More importantly, the observation that the pKa of such a monolayer is independent of its thickness indicates that the dielectric properties of the SAM/electrolyte interface are not a function of this thickness. This means, for example, that, in the absence of image charge effects49 the outer-sphere reorganization energy of a surface-attached redox couple20,21 should not change. Conclusions The principal conclusion that can be drawn from the present study is that, at least in the range of thickness defined by monolayers made from DBAD and MUA, the pKa of the monolayer is independent of its thickness. This (45) Zhang, C.-J.; Porter, M. D. J. Am. Chem. Soc. 1994, 116, 11616. (46) Chidsey, C. E. D.; Liu, G.-Y.; Rowntree, P.; Scoles, G. J. L. J. Chem. Phys. 1989, 91, 4421. (47) Hu, K.; Chai, Z.; Whitesell, J. K.; Bard, A. J. Langmuir 1999, 15, 3343. (48) Wang, J.; Bard, A. J. J. Phys. Chem. B 2001, 105, 5217. (49) Liu, Y.-P.; Newton, M. D. J. Phys. Chem. 1994, 98, 7162.

ω-Mercaptoalkanoic Acid Monolayers on Gold

result implies that the properties of the microenvironment of the SAM/electrolyte interface are also independent of monolayer thickness. It should also be noted that, as with MUA, the pK∞ of DBAD and MHA monolayers is considerably less than the pKa’s of alkane carboxylic acids in high ionic strength aqueous solutions (for example, the pKa of acetic acid in 2.5 M NaClO4 is 4.9450). This difference between the pK∞ measured in the present study as well as that reported in ref 31 and the solution pKa’s of alkane carboxylic acids at high ionic strengths is opposite to that predicted by the Born equation if the dielectric constant of the monolayer/electrolyte interface is less than that of the electrolyte solution.7 This observation, therefore, indicates31 that the microenvironment at the monolayer/ electrolyte solution interface either stabilizes the carboxylate anion and/or makes the carboxyl proton more acidic. Furthermore, if the charges are discreet,51,52 then the stabilization of charges produced at a monolayer/electrolyte solution interface due to image charge effects increases as the monolayer thickness decreases.49 pKa (and pK∞) should, therefore, decrease as the monolayer thickness decreases. However, calculations based upon the threephase (i.e., electrode/monolayer/solution) model proposed by Liu and Newton49 demonstrate that the size of the change in pKa effected by this change in image charge stabliziation in going from monolayers composed of MUA to those composed of DBAD is within the experimental error reported for the pKa’s in Table 1. This also means that image charge stabilization of the interfacial carboxylate anion cannot be responsible for the difference between the measured values of pK∞ (Table 1) and the pKa of an alkane carboxylic acid in aqueous solutions. (50) Rossotti, F. J. C.; Rossotti, H. The Determination of Stability Constants and Other Equilibrium Constants in Solution; McGraw-Hill: New York, 1961; p 21. (51) (a) Fawcett, W. R. J. Electroanal. Chem. 1994, 378, 117. (b) Fawcett, W. R.; Fedurco, M.; Kova´cˇova´, Z. Langmuir 1994, 10, 2403. (52) (a) Barlow, C. A., Jr.; MacDonald, J. R. J. Chem. Phys. 1963, 40, 1535. (b) Barlow, C. A., Jr.; MacDonald, J. R. J. Chem. Phys. 1965, 43, 2575. (c) MacDonald, J. R.; Barlow, C. A., Jr. J. Electrochem. Soc. 1966, 10, 978.

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It has already been suggested31 that the known43,53 difference in the stability of the syn (O-H and CdO on the same side of C-O) and anti (O-H and CdO on opposite sides of C-O) conformations of the carboxyl moiety could make the carboxyl proton more acidic. The syn conformation is the more stable form (with a stability constant in aqueous solutions of 104),54 so if the microenvironment at the monolayer/electrolyte interface forces the carboxyl moiety into the anti conformation, the SAM carboxyl proton should be more acidic than that associated with an alkane carboxylic acid in an aqueous solution. However, the confirmation of this suggestion must await better55 models for the detailed (molecular) structure of the monolayer/electrolyte interface than those that currently exist. Nevertheless, the thermodynamic results reported here are a probe of the microenvironment of this interface and, therefore, provide part of the information required for the construction of these models. Acknowledgment. I gratefully acknowledge the support of the U.S. Department of Energy, contract no. DEAC02-98CH10886. I also thank Professor Edmond F. Bowden (North Carolina State University) and Professor Harry O. Finklea (West Virginia University) for the gifts of the dibutanoic acid disulfide and the 6-mercaptohexanoic acid used in the present study. The many helpful discusssions with Dr. Stephen W. Feldberg (Materials Science Department, Brookhaven National Laboratory) and Marshall D. Newton (Chemistry Department, Brookhaven National Laboratory) are also greatly appreciated. I also thank Professor Christopher E. D. Chidsey and Dr. Hadley D. Sikes (Chemistry Department, Stanford University) for making the gold film electrodes used in this work. LA0348968 (53) (a) Peterson, M. R.; Csizmadia, I. G. J. Am. Chem. Soc. 1979, 101, 1076. (b) Meyer, R.; Ha, T.-K.; Frei, H.; Gu¨nthard, Hs. H. Chem. Phys. 1975, 9, 383. (54) Gandour, R. D. Bioorg. Chem. 1981, 10, 169. (55) Creager, S. E.; Rowe, G. K. J. Electroanal. Chem. 1997, 420, 291.