877
J. Phys. Chem. 1991, 95, 877-886
In the 180-exchanged reaction, the O2atoms in the internal SiOH groups or in the nonintact Si-0-Si groups in the nest type defect sites attained the equilibrium state at ca. 15 h. Moreover, the bands at 3505 cm-I were observed less in the IR spectra for the HZSM-5 having fewer defect sites. From these results, it is suggested that the broad band at 3505 cm-I can be identified with the SiOH groups of the hydroxyl nests. If aluminum atoms are inserted into hydroxyl nests in the HZSM-5 by alumination with AICI,, the band at 3505 cm-I will disappear. The IR spectrum for aluminated HZSM-5 prepared from the parent HZSM-5 (sample 9 in Table I) is depicted in Figure 8E. The aluminated HZSM-5 exhibited a new band at 3610 cm-’ attributed to the framework Si(0H)AI groups,23while the band at 3505 cm-’ disappeared. These results indicate that ~
~~~
(23) Jacobs, P. A.; Von Ballmoos, R. J . Phys. Chem. 1982, 86, 3050.
aluminum atoms can be inserted into hydroxyl nests in the highly siliceous HZSM-5 zeolite. In addition, as depicted in Figure 8, the band at 3740 cm-’ was dramatically decreased by the alumination. The fact that the band at 3740 cm-l decreases indicates that the external SiOH groups corresponding to the band at 3740 cm-l are consumed by the reaction of the SiOH groups with AICI,. It has been reported that aluminum atoms are introduced not only into the framework sites but also into extraframework sites as six-coordinate species by the alumination with AICI,. The six-coordinated extraframework aluminum species can be generated by the reaction of the SiOH groups of the band at 3740 cm-’ with AICI,.
Acknowledgment. We thank Prof. T. Takaishi and Dr. A. Endoh for suggesting the measurement of the amount of oxygen on defect sites by 180-exchange reaction and Mr. H. Shouji for 27AlMAS N M R measurements and for stimulating discussions. ’
Adsorbed *Hydroxy Thiol Monolayers on Goid Electrodes: Evidence for Electron Tunneling to Redox Species in Solution Cary Miller,* Pierre Cuendet, and Michael Gratzel Institut de Chimie Physique, Ecole Polytechnique Fgdgrale, CH- 1015 Lausanne, Switzerland (Received: June 14, 1990)
Au electrodes are derivatized by self-assembled organic monolayers of w-hydroxy thiols from aqueous solutions. The capacitance, ellipsometric, XPS, heterogeneous electron-transfer properties, and stability of these derivatized electrodes are consistent with the formation of a pinhole-free hydroxy thiol monolayer. In particular, the logarithmic dependence of the heterogeneous electron-transfer rate constant of Fe(CN)63- and Fe3+ in 0.1 M KCI on the thickness of the monolayer film and the lack of a temperature dependence at high overpotentials are consistent with electron tunneling through the full thickness of the self-assembledfilm. A tunneling constant fl of 0.9 per methylene group is measured. Tafel plots for Fe(CN),&/* and Fe2+I3+ couples display curvature at high overpotentials, which is consistent with Marcus electron-transfer theory.
Introduction The spontaneous self-assembly of amphiphiles provides a convenient technique by which compact monolayers with controlled thickness and structure can be prepared.I4 An interesting use of such self-assembled films is as insulating barriers at electrode surfaces. For monolayer films sufficiently free from pinhole defects, where redox species are effectively blocked from the electrode surface, the dominant mechanism of electron transfer should be by electron tunneling through the monolayer film. Hence, the rate of redox reactions at electrodes coated by such films should be strongly dependent on the structure and thickness of the insulating m o n ~ l a y e r . ~Electron tunneling through selfassembled monolayer films has been demonstrated in measurements by Polymeropoulos and Sagiv for alkanoic acid monolayers on AI/AI2O3/monolayer/AI capacitor structures6and by Li and Weaver for a series of cobalt complexes with thiol ligands assembled on Au electrodes.’ Attempts to measure electron tunneling through alkylsiloxy and alkanethiol monolayers to redox couples in solution have not been as successful.&’o In these studies the presence of defects within the monolayer films hampered the observation of electron tunneling through these electrode films. It is clear that the development of amphiphiles and adsorption conditions which can produce monolayer films of superior stability and perfection are required to observe this long-range electron transfer. The formation and structure of self-assembled monolayers has been the subject of many studies.14 The ability of an amphiphile * To whom correspondence should be addressed. Current address: De-
partment of Chemistry, University of Maryland, College Park, MD 20742.
to form a compact self-assembled monolayer and its subsequent stability at the surface depend on three principal interactions: the strength of binding of the headgroup for the surface, the attractive forces between the adsorbed amphiphiles within the monolayer, and the interaction of the amphiphiles’ terminal function with the surrounding ambient. Therefore, in order to produce as stable a monolayer as is possible, one should try to maximize each of these interactions. In particular, the choice of the terminal function should play an important role in determining the monolayer stability and packing. Previous work has shown that organosulfur compounds with a range of terminal functions can self-assemble onto Au surface^.^*"*'^ The packing of the amphiphiles on Au surfaces and their wetting characteristics as a function of the ( I ) Bigelow, W. C.; Pickett, D. L.; Zisman, W. A. J . Colloid Sci. 1946, I , 513. (2) Sagiv, J. J . Am. Chem. SOC.1980, 92, 102. (3) Allara, D. L.;Nuzzo, R. G. Langmuir 1985, I, 45. (4) Bain, C. D.; Troughton, E. B.; Tao, Yu-Tai; Evall, J.; Whitesides, G. M.; Nuzzo, R. G. J. Am. Chem. SOC.1989, 1 1 1 , 321. (5) Bockris, J. OM.; Khan, S.U.M. Quanfum Electrochemistry; Plenum Press: New York, 1979; Chapter 8, p 235. (6) Polymeropoulos, E. E.;Sagiv, J. J . Chem. Phys. 1978, 69, 1836. (7) Li, T. T.-T.; .Weaver, M. J. J. Am. Chem. Soc. 1984, 106, 6107. (8) Finklea, H. 0.; Robinson, L. R.; Blackburn, A.; Richter, B.; Allara, D.; Bright, T. Langmuir 1986, 2, 239. (9) Finklea, H. 0.; Avery, S.; Lynch, M.; Furtsch, T. Langmuir 1987, 3, 409.
(IO) Porter, M. D.; Bright, T. B.; Allara, D. L.; Chidsey, C. E. D. J . Am. Chem. Soc. 1987, 109, 3559.
(I I ) Nuzzo, R. G.; Fusco, F. A.; Allara, D. L. J. Am. Chem. SOC.1987, 109. 2358.
(12) Fabianowski, W.; Coyle, L. C.; Weber, B. A,; Granata. R. D.; Castner, D. G.;Sadownik, A,; Regen, S.L. Langmuir 1989, 5, 35.
0022-3654/91/2095-0877%02.50/0 0 1991 American Chemical Society
878 The Journal of Physical Chemistry, Vol. 95, No. 2, 1991
terminal function have received considerable attention!J3-'s In this work we report that alkanethiols with terminal hydroxyl functions adsorb onto Au electrodes from aqueous solution, forming remarkably compact and stable monolayers. The heterogeneous electron-transfer kinetics of redox couples are slowed at these hydroxy thiol derivatized electrodes. Quantitative analysis of the kinetics of the electron transfer is compared with expectations for tunneling through the electrode film.
Experimental Section Syntheses. w-Hydroxy thiols (HO(CH2),$H) were synthesized from the Corresponding w-hydroxyalkyl bromides or chlorides via the nucleophilic displacement of the halide with an excess of thiourea in 90% CH3CH20H:H20.I6 After the addition of 1.2 equiv of NaOH and refluxing for 30 min (causing the base hydrolysis of the thiouronium salt to give the mercaptan), the ethanol was removed with a rotary evaporator at reduced pressure, H 2 0 was added, and the crude product was extracted from the aqueous phase with diethyl ether. The resulting product was purified by column chromatography on silica gel (Fluka for chromatography, 60 mesh) with HCCI, as the mobile phase. 6-Chlore 1-hexanol (practical grade) and 1 1-bromcF 1-undecanol (purum) were purchased from Fluka and used as received. The HO(CH2),,Br ( n = 7, 8 , 9, IO, 12) were synthesized from the corresponding dibromides (Fluka, purum) by using the following procedure. One gram of the dibromide was dissolved in 30 mL of ethanol/H20 (80:20). The dibromide solution was heated to reflux, and 2 g of KOH was added. The mixture was allowed to reflux for 2 h at which time the solution volume was reduced by 50% under reduced pressure in a rotary evaporator. The resulting mixture was extracted with hexane to remove the unreacted dibromide. The remaining ethanol was removed via the rotary evaporator, and the resulting aqueous solution was extracted with diethyl ether, which was washed with small portions of H 2 0 to remove the excess KOH. Evaporation of the ether afforded the crude hydroxyalkyl bromide which was used without further purification. The HO(CH2),Br (n = 14, 16) were synthesized from the corresponding dihydroxyalkanes (Aldrich). A 0.2-g sample of the dihydroxyalkane compound was dissolved in 30 mL of acetone and brought to reflux. A 0.4-g sample of 4.8 M HBr (in H 2 0 ) was added and the mixture refluxed for 2 h. At this time the solvent was removed via a rotary evaporator under reduced pressure (removing the acetone and HBr), and the resulting slurry was then treated with thiourea without isolating the hydroxyalkyl bromide. The yields of these syntheses were low (1O-30% isolated yield when corrected for the recovered dibromide or dihydroxy starting material), especially for the shorter hydroxy thiols due primarily to the incomplete extraction of the product by diethyl ether. Elemental analysis was performed on the HO(CH2)IISH (C: 64.7% expected, 64.6% found; H: 1 1.8% expected, 12.0% found; S: 15.7% expected, 15.6% found; Br: expected, O.O%, found -and 15 mM Fel+ in 0.1 M KCI, 10 mM pH 7.4 Tris buffer and in 0.1 M KCI, IO mM HCI, respectively. The lines drawn through the points correspond to the best-fit lines. The results of the linear regressions are given in Table 11.
ammetric experiments were performed in the 0.1 M KCI supporting electrolyte to determine the voltage range in which the hydroxy thiol layers are stable. As the alkyl chain length of the hydroxy thiol is increased, the anodic and cathodic potential limits (the potentials at which the background currents rise steeply beyond ca. 20 MA/") are observed to shift to more positive and negative potentials, respectively. The anodic limit shifts from ca. 700 to 1 I50 mV and the cathodic limit from ca. -800 to -1 150 mV as the length of the hydroxy thiol increases from the hexyl to the hexadecyl derivative. Voltage excursions beyond these limits, which correspond to the potential required for the oxidation of the Au surface and the evolution of hydrogen, respectively, were observed to damage the monolayers, resulting in higher capacitances and more reversible voltammetry of the Fe(CN),4-/3couple. For electrodes derivatized by the shorter hydroxy thiols (hexyl to nonyl), both the electrode capacitance and the rate of redox kinetics were observed to increase slowly with continuous cycling between the anodic and cathodic limits. Such instability is not observed with electrodes derivatized with the longer hydroxythiols, which can be cycled hundreds of times between the anodic and cathodic limit with no significant change in either their capacitances or kinetics of heterogeneous electron transfer to species in the electrolyte. The reduction kinetics of Fe3+ and Fe(CN)63- in 0.1 M KCI (IO mM HCI and IO mM pH 7.4 Tris, respectively) at the hydroxy thiol derivatized electrodes were assessed from the initial portion of cyclic voltammograms. Plots of the logarithm of the measured apparent electron-transfer rate constant versus the electrode potential (Tafel plots) for the hydroxy thiol derivatized electrodes are shown in Figure 5. In these plots only the data in the
w-Hydroxy Thiol Monolayers on Au Electrodes
The Journal of Physical Chemistry, Vol. 95, No. 2, 1991 881
A
TABLE 11: Linear Remession Results from Data in Firmre 3 Fe3+ hydrocarbon Fe(CN)tchain length' interced slow . (ale . , interceptb slow . (aM 6 -3.05 11.5 (0.68) -6.06 6.28 (0.37) ~
-
1
7 8 9
-3.68 -4.43 -4.95 -5.44 -5.74 -6.05 -6.81 -6.90
IO II I2 14
16
8.34(0.49) 6.75 (0.40) 5.91 (0.35) 6.05 (0.36) 5.87 (0.35) 4.91 (0.29) 4.25 (0.25) 3.38 (0.20)
-6.37 -6.71 -7.03 -7.17 -7.37 -7.40 -7.92 -8.40
T
,
n
5.79 (0.34) 5.41 (0.32) 5.08 (0.30) 4.95 (0.29) 4.62 (0.27) 4.09 (0.24) 3.84 (0.23) 3.64 (0.22)
'The number of methylene groups ( n ) in the hydroxy thiol with a formula HO(CH,)SH. blog (k,,/(cm/s)). 'The Tafel slopes and the electron-transfer symmetry factor (Y (in parentheses). E (Volt)
B 7
4 8
4 4
4 2
4 0
48
Q8
Q4
QZ
00
02
04
08
-
No added C1
08
+0.6
+0.5
0.0
-0.2
overpotential / V
Figure 6. Tafel plots obtained in 15 mM solutions of Fe(cN)6+/h (open symbols) and Fez+/)+(solid symbols) in 0.1 M KCI, IO mM pH 7.4 Tris buffer and 0.1 M KCI, IO m M HCI, respectively, at an Au electrode derivatized with a HO(CH2),6SHmonolayer.
low-overpotential %near region" were used. For all these data, the diffusion-limited rate of mass transfer of the redox species to the electrode measured at a bare electrode was at least 20 times higher than the maximum measured rate so that diffusion limitations were avoided. Increasing or decreasing the scan rate of these voltammetric experiments by a factor of 5 had.no effect on the measured rates. The series of curves for both Fe3+ and Fe(CN),'- are reasonably linear with some curvature observed at larger overpotentials. The line through each set of data corresponds to the best fit line; the results of these linear regressions are given in Table 11. It is observed that, with increasing length of the hydroxy thiol, both the intercept and slope of the Tafel plots decrease. In the course of measuring the reduction kinetics of Fe3+ and Fe(CN)63-, significant deviation from Tafel behavior at large overpotentials was observed. In order to investigate the sources of these deviations, the reduction and oxidation kinetics for both Fe2+/3+and Fe(CN)64-/3-were measured at as broad a voltage range as was possible. Figure 6 shows two sets of Tafel plots for these redox couples obtained with a HO(CH2)16SHderivatized electrode. The Tafel plot for the Fe(CN),4-/3- couple displays much more curvature than the one for the Fe2+l3+. In an effort to probe the level of "pinhole" defects within these hydroxy thiol monolayers, we investigated the effect of small concentrations of CI-on the reduction kinetics of Fe3+ at a HO(CH,),SH treated Au electrode. This method relies on the known dependence of halide adsorption on the kinetics of the aquo iron coupleam Figure 7 shows two sets of voltammetric curves for the reduction of 5.0 mM Fe(C104)3before and after the addition of 2.0 mM KCI to the electrolyte. For the bare Au electrode (Figure 7A), a significant improvement in the kinetics of the Fe3+reduction are observed with the addition of the 2 mM CI-. However, as ~
~
~~~
(20) Webcr, J.; Samec, Z.; Marecek, V. J. Electroanal. Chem. 1978.89.
21 I .
E (Volt) Figure 7. Effect of CI-on Fe3+reduction kinetics. Two sets of voltammograms obtained in 5.0 mM Fe3+,50 mM HCIO,, prepared with purified Fe(CIO& in the absence and presence of 2 mM KCI as indicated in the figure. The voltammograms in (A) were obtained by use of a bare Au electrode, while those in (B) were obtained by using a Au electrode derivatized with a HO(CH&SH monolayer. The scan rate used for these voltammograms was 0.5 V/s.
seen in Figure 7B, the Fe3+reduction kinetics at the HO(CH2)&H treated electrode are not observed to change significantly. Similar experiments using electrodes coated with the undecyl and hexadecyl hydroxy thiols showed the same lack of a strong dependence on the presence of chloride even when the chloride concentration was 0.3 M. Some variability in the electrode kinetics of electrodes treated with the same hydroxy thiol was observed. In some cases the cyclic voltammograms displayed sigmoidal waves at lower overpotentials or shifts in the voltammetric wave to lower overpotentials. Such variability appears to be due to variations in the cleanliness of the Au electrodes or the hydroxythiol solutions. In these cases rinsing the electrode sequentially with H 2 0 , ethanol, H 2 0 and then repeating the exposure of the electrode to the hydroxy thiol solution was found to eliminate these effects so that the current at a given overpotential could be reproduced to within 30%. In order to probe the activation characteristics of electrontransfer reaction at these blocked electrodes, the temperature dependence of the Fe(CN)$-l3- redox kinetics was measured on a Au electrode derivatized with HO(CH2)16SH.Figure 8 shows a plot of the apparent electron-transfer rate constant as a function of the electrode overpotential obtained at the temperatures shown in the figure. For each temperature the Fe(CN)64-/3- redox potential was measured with a bare Au electrode. Within the temperature range 0.6-71.2 OC, the redox kinetics in the voltage range shown are not a strong function of the temperature. This experiment was repeated several times with electrodes treated with either undecyl or hexadecyl hydroxy thiols with the same lack of a temperature dependence on the observed electrode kinetics at high overpotentials (v > 0.4 V). At low overpotential values (-0.1
882 The Journal of Physical Chemistry, Vol. 95, No. 2, 1991
Figure 8. Plots of the apparent heterogeneous electron-transferrate kapp versus electrode overpotential for the reduction and oxidation of a 15 m M Fe(CN)$-/)- in 0.1 M KCI, IO m M pH 7.4 Tris buffer electrolyte at a Au electrode derivatized with a HO(CH2),6SHmonolayer. The curves correspond to data taken at different temperatures as indicated in the figure, and the positive and negative direction on the ordinate corresponds to the reductive and oxidative electron transfer, respectively. T
500 nA
I
E (Volt)
Figure 9. Two voltammograms obtained at a HO(CH2),,SHderivatized Au electrode in a solution containing 1.7 mM Fe(CN)6t/)- in 0.1 M KCI, IO m M pH 7.2 Tris buffer electrolyte. The solid curve shows the voltammogram bcfore the exposure to sulfochromic acid while the dotted curve shows the voltammogram at the same electrode after a 15-s exposure to sulfochromic acid at 85 OC. The scan rate for these voltammograms was 200 mV/s.
to 0.3 V), some increase in the observed heterogeneous electron-transfer rate was observed with increasing electrolyte temperature. Experiments of the temperature dependence of electrodes coated with the shorter hydroxy thiols could not be done, owing to their poorer stability under continuous electrochemical cycling. The stability of the HO(CH,),SH treated electrodes to dissolution by solvent and chemical attack was also investigated. It was found that the hydroxy thiol monolayers show impressive stability to hot solvents and even sulfochromic acid treatments. For example, a HO(CHJI1SH treated Au electrode was exposed to ethanol at reflux for 2 min and then to 2 min to sulfochromic acid at room temperature for 2 min with no change in its capacitance. A 15-s exposure of this electrode in sulfochromic acid at 85 OC resulted in only a 20% increase in the electrode capacitance. Figure 9 shows cyclic voltammograms obtained in a 1.7 mM Fe(CN)d-, 0.1 M KCI solution obtained before and after the 85 OC sulfochromic acid treatment. While one can see several shoulders on the voltammetric wave obtained after the hot sulfochromic acid treatment, no drastic change in the Fe(CN)63reduction kinetics is observed.
Discussion The spontaneous self-assembly of hydroxy thiols from aqueous solution onto Au electrodes produces monolayers with exceptional insulating properties. In our experience and by comparison with other results, the hydroxy thiols form more stable and blocking monolayers than equivalent alkanethiols. We postulate that the hydroxy terminal function stabilizes the monolayer by hydrogen-bonding to adjacent hydroxyl groups and to the aqueous solution. This should increase the attractive forces within the monolayer and decrease the monolayer/aqueous surface energy when compared to a methyl terminal function. In addition, the stability and the lack of measurable defects of these monolayers allow one to put an insulating barrier between the electrode surface
Miller et al. and redox species in solution. The mechanism for electron transfer at these electrodes appears to be dominated by electron tunneling through the hydroxy thiol layer. The focus of the following discussion is to justify these assertions and to present the significance of this work. The capacitance data presented in Table I and Figure 1 support our assertion that the hydroxy thiols form compact monolayers at Au electrodes. The relative dielectric constant of 3.0 calculated from these data compares favorably to the values reported for alkanethiol (2.6)'O and alkanoic acid monolayers (2.9)." For these systems, IR spectroscopy has shown that the monolayers adsorb in an "all-trans" config~ration.~J~J' So it is reasonable to assume that the hydroxythiols form similar closed-packed structures. However, none of the measurements presented here allows us to determine the exact structure or degree of crystallinity of these films. The calculation of the relative dielectric constant will likely be an upper bound of the true value due to several assumptions made. First, it was assumed that the surface area of the electrode is identical with the geometric area. Because a certain roughness is expected at these electrodes," the true area of the electrode is likely higher, causing an overestimation of the relative dielectric constant. In addition, spectroscopic investigations of alkanethiol monolayers adsorbed on Au surfaces have indicated that the monolayer adsorbs at a significant angle to the surface normal." This tilting of the adsorbed monolayer if present in our system would result in each methylene group of the hydroxy thiols adding less than 1.25 A to the monolayer thickness, again resulting in an overestimation of the dielectric constant. The linearity of the reciprocal capacitance plot of Figure 1 extends to the shorter hydroxy thiols (6 I n I 9), suggesting that the level of organization (degree of solvent and ion blocking) is similar to the longer hydroxy thiols.2' In the study of alkanethiols adsorbed at Au electrodes by Porter et a1.,I0 it was observed that the plot of 1/C versus the hydrocarbon chain length curved for chain lengths of fewer than nine carbons. The authors ascribe this curvature, observed in KCI electrolytes, to the ability of CIions to penetrate the thinner monolayers. That we do not observe this effect suggests that the hydroxy thiols form tighter structures than the corresponding alkanethiols. Furthermore, the insensitivity of the Fe3+ reduction kinetics to additions of CI- to the electrolyte also indicates a lack of CI- permeability through these layers. It must be pointed out that, in the study of Porter et al.,IOthe capacitances were measured at much lower scan rates than in this study. For both the hexyl and ethyl hydroxy thiol derivatized electrodes, the electrode capacitances were measured at 0.5 V/s and were found to be within 10% of that measured at 51.2 V/s. These two hydroxy thiols were chosen because they produce electrodes having the largest capacitance, allowing a semiquantitative measurement of the electrode capacitance at these low scan rates. Unfortunately, we were limited to these scan rates due to the sensitivity of our current measurements and the small size of our electrodes. Not only the presence of the polar terminal function but also the adsorption of the amphiphile from aqueous solutions seems important in obtaining optimally packed monolayer films. We repeated the derivatization procedure of Bain et al! in which the Au surfaces are exposed to ethanolic solutions of the hydroxy thiols. While the capacitances of electrodes derivatized by a 30-min exposure to the ethanolic hydroxy thiol solution do indicate substantial adsorption of thiol, they are roughly twice that observed as compared to the capacitances of electrodes derivatized with aqueous solutions. Rinsing with ethanol, H 2 0 , and ethanol followed by repeated 30-min exposure to the ethanolic hydroxy thiol solution caused no change in the electrodes' capacitances. Thus, it seems that the level of adsorption observed after the first exposure corresponds to the limiting coverage in this solvent. Because a subsequent treatment of these electrodes in an aqueous hydroxy (21) Such a claim for the shortest hydroxy thiol ( n = 2 ) becomes less tenable given that the double-layer capacitance is neglected in the analysis. The capacitance of the mercaptoethanol treated electrode is only a factor of 3 lower than that of the bare electrode.
w-Hydroxy Thiol Monolayers on Au Electrodes thiol solution lowered their capacitance to others derivatized with aqueous solutions, a fault in the cleanliness of the Au electrodes cannot be used to explain the lower absorption from ethanolic solutions. However, because our electrode pretreatment and hydroxy thiol exposure conditions are not precisely identical with those of Bain et al., it is still possible that our failure to obtain tight packing of these thiols from ethanolic solution is a result of these small differences. We interpret the difference in the packing of the hydroxy thiols self-assembled from ethanolic and aqueous solutions as resulting from two effects. For the aqueous solution, the solution of the hydroxy thiol is saturated. Thus, no free energy is lost by 'precipitating" the amphiphile in a self-assembled monolayer at the Au surface. Secondly, the ethanol, being more able to solvate the amphiphile, may be incorporated within the monolayer hindering its full packing. This effect of solvent incorporation was also reported in the study of Bain et a1.4 in which the solvent dependence on the ellipsometric thickness and the surface hydrophobicity (as measured by H20and hexadecane contact angles) of self-assembled hexadecanethiol monolayers on Au surfaces was compared to the ability of the solvent to be incorporated in the monolayer. The ellipsometric determination of the adsorbed monolayer thicknesses shown in Figure 2 and the XPS experiments resulting in Figure 3 were performed to add further weight to our claim that the absorption of the hydroxy thiols produce single monomolecular films. In agreement with the capacitance measurements, we see a linear increase in the ellipsometric thickness (as measured by a change in the ellipsometric angle A) as the length of the hydroxy thiol is increased. The point corresponding to 18 methylene units corresponds to the thickness of an Au surface derivatized with octadecyl mercaptan from ethanolic solution. This point was included as an internal control of the thickness determination. The slope of the best fit line shown in Figure 2 corresponds to a 1.5-A increase in the monolayer thickness with each addition of a methylene group. This 1.5 A per methylene slope is identical with that reported for a series of alkanethiols adsorbed on Au surface^.^*'^ In Figure 3 are shown the change in the integrated C Is signal and the Au 4f signal with increasing length of the hydroxy thiol used to form the monolayer film. Because of the possibility of the adventitious adsorption of carbonaceous contaminants on the hydroxy thiol treated Au surface, the XPS data cannot be used to quantitatively measure the thickness of the adsorption. However certain trends are significant. As the length of the hydroxy thiol is increased, we seen an approximately linear increase in the amount of carbon adsorbed on the Au surface and an increasing attenuation of the Au signal by the monolayer film. Both these trends are consistent with a regularly increasing thickness of the monolayer film. By comparison with the octadecyl mercaptan point also shown in Figure 3, one can assert that the level of the carbon signal and the attenuation of the Au signal are consistent with a single monolayer film. The formation of multilayer films of the hydroxy thiol monolayers does not occur under the experimental conditions used here. As the length of the hydroxy thiol increases, we observe a gradual attenuation of the S 2p signal and an approximately constant 0 Is signal. Such observations are consistent with the hydroxy thiols being oriented with their thiol group toward the Au electrode exposing the hydroxyl function. Faradaic Processes. The electron-transfer kinetics at these hydroxy thiol derivatized electrodes are also consistent with the formation of compact low defect density monolayers. As observed in Figure 4, the reduction kinetics of Fe(CN)be/* couple depend strongly on the length of the hydroxy thiol. With each increase in the number of methylene groups of the hydroxy thiol, the thickness of the insulating barrier increases, reducing the rate of electron transfer across the monolayer. As seen in Figure 4, the voltammetric waves are stable to repeated cycling of the electrode potential. Quantitative analysis of this barrier to electron transfer can be made by constructing Tafel plots for the reduction as a function of the number of carbon atoms in the hydroxy thiol which are shown in Figure 5. I t is expected that if electron tunneling
The Journal of Physical Chemistry, Vol. 95, No. 2, 1991 883
A 'I
P
I
-3i
m
-'i -8
-7
,
, 7
,
(
,
9
: ; , ; , ; / 11
13
15
17
Hydrocarbon Chain Length
B
7J 5
"
7
"
9
.
'
.
11
"
13
' 15
-. 17
Hydrocarbon Chain Length Figure 10. Plots of the log k, versus the number of methylene groups, n, in the hydroxy thiol, HO(&l2)$H, at the derivatized Au electrodes ('tunneling plots"). The k,, values were obtained through the extrapolation of the Tafel plots in Figure 3 at - 0 . 2 3 4 and - 0 . 5 2 2 V for the Fe(CN);- and Fe3+ curves, respectively. The line drawn in (B) for the Fe3+ data corresponds to the best-fit line and is plotted to allow one to observe the curvature in this plot.
across the monolayer is implicated in the electron transfer, the reduction current at any potential should decrease exponentially with the monolayer thickness according to the expressionZZ
i = ioe-Bd
(2)
where io is the observed current in the absence of the insulating barrier, d is the thickness of the electrode film, and j3 is the electron tunneling constant. For a rectangular barrier this j3 gives a measure of the barrier height though the equationZ3
(3) where m is the free electron mass, h is Planck's constant, and V is the height of the barrier in electronvolts. Substituting in the values of the constants in eq 3, one obtains = 1.025Vi/2
(4)
where the barrier thickness is in angstroms and the barrier height is in electronvolts. From the Tafel data for both the Fe(CN),3and Fe3+ reductions, the apparent reduction rate constants can be extrapolated to a common potential which can be analyzed according to this tunneling expression. Figure 10 shows the dependence of log k, (the apparent rate constant for the reduction of Fe(CN)6' and Fey+) at -0,234 and -0.522 V,extrapolated from the data in Figure 5. These potentials (which correspond to 0.0 V vs Ag/AgCl, saturated KCI) were chosen to minimize the extrapolation required from the Tafel plots. The plot obtained for Fe(CN)63- reduction (Figure 8A) displays marked curvature from a linear change of the log kaPpwith the length of the hydroxy (22) Marcus, R. A,; Sutin, N. Eiochim. Eiophys. Acrrr 1985,811, 265. (23) Hartman, T.E.J. Appl. Phys. 1964, 35, 3283.
884
The Journal of Physical Chemistry, Vol. 95, No. 2, 1991
thiol while that for the reduction of Fe3+ shows much less curvature. The reason for the differences in these plots stems from the difference in the Tafel plots for these two couples. As seen in Table 11, the Tafel slopes for both reductions decrease with increasing length of the hydroxy thiol monolayer. The decrease in the Tafel slope between the hexyl and hexydecyl hydroxy thiol is over a factor of 3 for the ferricyanide data while less than a factor of 2 for the ferric ion data. As the hydroxy thiol layer becomes thicker, the kinetic data must be obtained at higher and higher overpotentials. This is basically a limitation on the sensitivity of the current measurement. If the tunneling of the electrons through the hydroxy thiol layer is elastic, all of the electron overpotential will be available for the electrochemical reduction. According to the Marcus theory for heterogeneous electron-transfer reaction^,^^^^^ it is expected that as the driving force of the reduction approaches the reorganization energy of the electron transfer, the Tafel slope should approach zero according to the equation22,2sq26
where n is the number of electrons transferred, F is Faraday’s constant, 7 is the electrode overpotential, and A is the reorganization energy of the couple. The Tafel plots for the two redox couples obtained at a large range of overpotentials (Figure 6 ) show just this sort of decreasing Tafel slope with increasing overpotentials. From the extent of curvature between the two sets of Tafel plots in Figure 6 , one can conclude that the reorganization energy for the Fe2+/)+couple is much larger than that for Fe(CN)64-/3-. This is of course in agreement with reported measurements for the reorganization energies of these As a consequence of the Fe2+/3+couple’s higher reorganization energy, the data for the ferric ion reduction should give a truer picture of the dependence of the electron-transfer rate on the thickness of the monolayer film. The near linearity of the plot in Figure 8B is supportive for electron tunneling through these monolayer covered films as the dominant mechanism of electron transfer. From the measurement of the tunneling constant (3, one can obtain a measure of the barrier height for the hydroxy thiol films. Because of the variation in the Tafel slope with the length of the hydroxy thiol layer, (3 cannot be accurately measured from rate constants extrapolated from Tafel plots. This is because the (3 would then depend on the potential at which one extrapolates the apparent rate constant. To avoid this problem of extrapolation, we have taken the ratios of the measured apparent rate constants at the same potentials for pairs of electrodes coated with hydroxy thiol layers differing by 1, 2, or (for the longest hydroxy thiols) 4 methylene groups in length. Figure 11 shows a plot of the calculated /3 values versus the average of the number of methylene units of the two hydroxy thiol layers used to calculate (3. If one measures the (3 value (which is ideally the slope of the plots in Figure IO) using the data from electrodes coated with hydroxy thiols very close in length, small deviations in the kinetics give rise to large fluctuations in the calculated (3. As seen in Figure 1 I , there is considerable scatter in the calculated (3 values which stems most likely from differences in the true areas of the electrodes and from small differences in the packing of the self-assembled hydroxy thiol layers. No attempt was made to correct for the true surface area, but this would likely decrease the scatter. An average j3 value of approximately 0.9 (per methylene group) can be obtained from these data. Even though considerably more curvature is present in the “tunneling plot” for the Fe(CN),3- than (24) Marcus, R. A . J . Chem. Phys. 1965, 43, 679. (25) Tyma, P. D.; Weaver, M. J . J . Electroanal. Chem. 1980, I l l , 195. (26) Gr&el, M. Heterogeneous Photochemical Electron Transfer; CRC Press: Boca Raton, FL, 1989; Chapter I , p 36. (27) Mollers. F.: Memming, R. Ber. Bunsen-Ges. Phys. Chem. 1972, 76, 415. (28) Morisaki, H.; Ono, H.; Yazawa, K. Proc.-Elecrrochem. Soc. (Photoelectrochem. Electrosynth., Semicond. Mater.) 1988, 88, 436.
14
Miller et al.
. 0
0
0
: 0 0
8 I 0
1
04’
5
I
12
14
15
Avcragc Hydrocnrbon Clmin Length
Figure 11. A plot of the calculated tunneling constant, @, calculated from the ratio of the reduction kapp(for Fe(CN)63- and Fe3+ a s indicated in the figure) measured at Au electrodes derivatized by hydroxy thiol monolayers whose lengths were different by I , 2, or 4 methylene groups. These @ values are plotted versus the average length of the hydroxy thiol monolayers used for their calculation. (See text.)
for Fe3+ (Figure IO, A vs B), the average (3 values obtained for either couple are the same. With the large scatter in the (3 determination shown in Figure 11, one should be quite concerned with the accuracy in taking the average value. Support for this average value can be obtained by comparing the reported standard rate constant for the Fe(CN),)- reduction on a bare Au electrode (0.031 ~ m / s with ) ~ ~ that obtained from the Tafel plot for the same reaction at the Au electrode derivatized with the hexadecyl hydroxy thiol ( 2 X IO-* cm/s). The ratio of the two rate constants should give a measure of the tunneling barrier. If one uses a (3 value of 0.9, the attenuation in the electron-transfer rate caused by the hexadecyl hydroxy thiol monolayer would be 5.5 X which is in close agreement with the ratio of the two standard rate constants (6.5 x 10-7). Using this (3 and assuming 1.25 A per methylene group, this corresponds to a barrier height of 0.5 eV. This is considerably smaller than the barrier measured by Polymeropolous and Sagiv for AI/AI2O3/alkanoic acid monolayer/Al structures (2.6 eV)6 or by Li and Weaver7 for a series of cobalt complexes with thiol ligands ( 2 eV). In contrast, the (3 value per methylene group found here agrees well with that reported for the distance dependence of the electron transfer of donor-accepter pairs separated by saturated, rigid hydrocarbon linkages ( 1 . 1 5 - - 0 . 9 8 ) . ~ ~In~ ~these works, an important through-bond coupling of the electron transfer along the a-bond network of the spacer group is postulated to give rise to the observed low (3 value. In a more closely related system, Auweraer et al. have reported a tunneling barrier height of 0.5 eV for the light sensitized hole injection across a hydrocarbon monolayer formed via the Langmuir-Blodgett technique.34 The large discrepancy between these values of (3 is not easily resolved. In all these systems, one would expect a similar coupling along the structurally similar hydrocarbon spacer groups. The coupling along a a-bonded hydrocarbon network is expected to be strongly dependent on the bond angles of the hydrocarbon chains.30 One could postulate that, in the work of Polyermopolous and Sagiv? the alkanoic acid monolayers between 77 and 4 K have frozen in an unfavorable geometry while the hydroxy thiols in this (29) Marecek, V.; Samec, Z.; Weber, J . J . Electroanal. Chem. 1978,94, 169. (30) Closs. G . L.; Calcaterra, L. T.; Green, N. J.; Penfield, K. W.: Miller. J . R. J . Phys. Chem. 1986, 90, 3673. (31) Beratan, D. N. J . Am. Chem. Soc. 1986, 108,4321. (32) Oevering, H.; Paddon-Row, M. N.; Heppener, M.; Oliver, A. M.; Cotsaris, E.; Verhoeven, J . W.;Hush, N. S . J . Am. Chem. Soc. 1987, 109, 3258. (33) Paddon-Row, M . N.; Oliver, A . M.; Warman, J . M.; Smit, K . J.; de Haas, M. P.; Oevering, H.; Verhoeven, J. W. J . Phys. Chem. 1988,92,6958. (34) Van der Auweraer, M.; Verschuere, 8 . ; Biesmans, G.; De Schryver, F. C.; Willig, F. Langmuir 1987, 3, 992. (35) Devault, D. Q.Reu. Biophys. 1980, 13, 387.
w-Hydroxy Thiol Monolayers on Au Electrodes study at room temperature can rotate and vibrate about their central axes, allowing intermittent optimal geometries for the coupling along the hydrocarbon chains. With increasing temperature the conductivity of the alkanoic acid monolayers was observed to increase, which the authors ascribed to a temperature-dependent "impurity conduction" but may be due in part to motion about the hydrocarbon chain axis. However, using the same line of argument, one would expect that in the work of Li and W e a ~ e r which ,~ was also done at ambient temperatures, thermal motion along the thiol ligand spacer would afford a similar chance for an optimal geometry. As the validity of these speculations is by no means certain, the discrepancy in the p values remains a vexing problem requiring further study and comparison. What we wish to stress here is that there are other experimental determinations of low tunneling barrier heights and that the presence of such a low barrier in this work does not require the abandonment of the electron tunneling mechanism. The temperature dependence of the Fe(CN)6e/3- redox kinetics at the HO(CH2),,SH and HO(CH2)$3H treated electrode is also consistent with the postulate of electron tunneling through the monolayer. If the monolayer does not undergo a drastic change in thickness, tunneling barrier height, or level of defects within a certain temperature range, one would expect the probability of electron tunneling through the insulating barrier to be independent of t e m p e r a t ~ r e .One ~ ~ would still expect some dependence of the Fe(CN),"/'- kinetics on temperature because, while the electron tunneling probability should be temperature independent, the reduction of ferricyanide would still be an activated process. This activation energy according to the Marcus theory approaches zero as the overpotential approaches the reorganization energy of the couple. At overpotentials above the reorganization energy of the redox couple, electrons and holes of higher energy than the Fermi level become involved in the electron transfer so that the electrode reaction remains barrierless. Thus, one would not expect to see the "inverted region" (predicted by the Marcus theory for homogeneous electron transfers where the activation energy increases with increasing driving force for the reaction above the reorganization energy) for heterogeneous electron transfers from metallic electrodes. So, at electrode potentials above the reorganization energy of the redox couple, the electron transfer should be temperature independent. The data shown in Figure 8 at large overpotentials above the reorganization energy of the Fe(CN)6e/3couple (ca. 0.4-0.7 eV27-29)do show only small differences with temperature. Thus over the temperature range studied, the insulating properties of the hydroxy thiol monolayers do not change significantly, in agreement with the electron tunneling mechani~m.~~ Defect Level Assessment. While both the dependence of the ferricyanide reduction kinetics on the monolayer thickness and its near temperature independence at high overpotentials are consistent with electron tunneling through the full thickness of the monolayer film, the possible presence of defects must be addressed. Defects can be of the form of small pinholes exposing the electrode surface or collapse sites9 in the monolayer where a lower packing density of the amphiphiles and the subsequent collapse of the alkyl chains cause the monolayer to be significantly thinner than the bulk monolayer film. Because of the strong dependence of the electron-transfer kinetics on the thickness of the insulating film, even a small coverage of such defects could be entirely responsible for the current measured. In order to assess the effect of such defects, we have used the theoretical results of Amatore et at., who have proposed a model for the redox kinetics of partially blocked electrode^.^^ According (36) The lack of a temperature dependence is admittedly at odds with our earlier postulate that thermal motion about hydrocarbon chains aids in the 'coupling" of the electron across the monolayer. The independence in the electron-transfer rate with temperature in the 0-70 OC range would suggest that whatever motion aids in the electron transfer is sufficiently fast at 0 "C to create a distribution of hydrocarbon geometries which does not vary significantly within this temperature range. (37) Amatore, C.;Sav€ant, J. M.; Tessier, D. J . Electroanal. Chem. 1983, 147. 39.
The Journal of Physical Chemistry, Vol. 95, No. 2, 1991 885 to this model, if an electrode has defects of either pinhole or collapse types which are separated by distances greater than the characteristic diffusion length of the experiment (roughly IO bm for the cyclic voltammetry used here), sigmoidal voltammetric waves characteristic of an array of microelectrodes should be observed. Such sigmoidal waves were observed in some electrode preparations but could be eliminated completely with repeated exposures to the aqueous hydroxy thiol solution. The electrodes used in this study did not display such sigmoidal waves, and so these types of widely spaced defects should not be of sufficient density to alter the measured kinetics. Electrodes covered by defects separated by distances smaller than the characteristic diffusion length would give rise to voltammetric curves indistinguishable in shape from those obtained at electrodes with continuous insulating layers. The possible presence of this type of defect giving rise to the observed results is therefore somewhat harder to discount. If the measured electrode kinetics for the hydroxy thiol treated electrodes was due to such closely spaced defects, one would not anticipate a linear correlation of the log k,,, at an arbitrary potential with the hydroxy thiol monolayer thickness. This is because the coverage of these defects and their standard rate constants would have to vary consistently with small changes in hydroxy thiol length. Most likely, the presence of this type of defect would cause increased variability in the measured electrode kinetics for a set of identically treated electrodes. SucR variability in electrode kinetics was reported for alkanethiol and alkylsiloxyl monolayers at Au electrodes.s-10 However, the good correlation between the length of the hydroxy thiol and the rate of the log kappand the reproducibility of the electrode kinetics from one electrode preparation to another are inconsistent for defect-dependent electron-transfer kinetics. The lack of a temperature dependence on the reduction rate for the Fe(CN)63- is also inconsistent for a diffusion-limited reaction to defects on the surface. Within the 70 O C range studied, the diffusion coefficient for the Fe(CN)?- ion would be expected to increase by 3-4 times.38 This should give rise to a 3-4 times increase in the measured rate constant at a given electrode potential unless the number of these defects or their size changed to compensate for the change in the diffusion coefficient with temperature. In addition, the kinetics for electron transfer are strongly dependent on the electrode potential, which would not be expected for a diffusion-limited reaction. Further indications about the possible presence of defects can be inferred from the Fe3+reduction kinetics at these hydroxy thiol coated electrodes. The insensitivity of the Fe3+reduction kinetics of an electrode coated with the hexyl hydroxy thiol to small concentrations of CI- is strong evidence against the presence of defects exposing the Au surface. The influences of small concentrations of specifically adsorbed anions on the redox kinetics of the Fe(H20),)+ complex have been well d o c ~ m e n t e d . ~The ~,~~ presence of as little as M CI- in the electrolyte can give rise to an increase in the heterogeneous electron-transfer rate constant observed at Au electrodes by a factor of 100. After purification of the Fe(C104)3salt to remove the CI- impurity, we could observe significant sluggishness of the reduction kinetics of the Fe3+couple at a bare Au electrode. With the addition of 2 mM KCI, the voltammogram becomes nearly reversible as can be seen in Figure 7A. The addition of the CI- results in at least a 40 times increase in the heterogeneous electron-transfer rate constant. This estimate of the increase in the reduction kinetics is extremely conservative; the current at the half-wave potential for the voltammogram obtained in the presence of CI- was compared to the current at the same potential in the absence of CI-. The voltammograms shown in Figure 7A serve to show that, for bare Au electrodes, the addition of CI- gives rise to a dramatic increase in the Fe3+ reduction kinetics. For the HO(CH,),SH derivatized electrode, a slight increase in the rate constant of perhaps 20%can be seen from the voltammograms shown in Figure 7B upon the addition (38) Mills, R.; Lobo, V. M. M. Seljdijfusion in Electrolyte Solutions; Physical Sciences Data 36; Elsevier: New York, 1989; Appendix I, p 313. Weber, J. J . Electround. Chem. 1977,77, 163. (39) Samec, Z.;
886 The Journal of Physical Chemistry, Vol. 95, No. 2, 1991 of 2 mM KCI to the electrolyte. If the reduction current measured at the derivatized electrode was due to closely spaced defects which expose microscopic areas of the Au electrode to the solution, one would expect at least a 40 times increase in the observed electron-transfer rate constant which was observed for the bare electrode. The nearly insignificant augmentation of 20% observed for the reduction rate constant of the HO(CH2)6SH treated electrode precludes the presence of these defects. Indeed, the size of this change in the reduction kinetics may be related to the exchange of CI- for H 2 0 in the inner coordination sphere of the Fe(H20)63+ion (association constants: pK, = 1.5, pK2 = 0.6, pK3 = -1.4, pK4 = -1 .9).40 For solutions prepared from FeCI, in 0.1 M KCI, typically we see an increase in the reduction rates by 50-80% as compared to solutions prepared from Fe(C10J3 in HCI04. The hexyl hydroxy thiol was used for this experiment because it produces the thinnest monolayers of the monolayers we have studied using redox kinetics measurements. Defects in the packing of this short hydroxy thiol would be more likely to expose bare Au surface than the longer hydroxy thiols. The presence of collapse sites in which the Au surface is covered (so as to preclude specific adsorption of Cl- but where the thickness of the adsorbed monolayer is significantly thinner that the bulk monolayer thickness) remains as the final type of defect possible for these hydroxy thiol derivatized electrodes. Their presence within such films on the polycrystalline Au electrodes seems almost certain. It is difficult to imagine how on these rough electrodes the monolayer could be perfectly uniform in thickness while conforming to the underlying Au surface. The underlying Au surface would thus cause some modulation in the thickness of the monolayer film. However, as stated before, the insensitivity of the redox kinetics to temperature and the reproducibility of the kinetics between electrode preparations argue against the dominant electron transfer occurring through a small number of defect sites of the hydroxy thiol monolayers. The presence of a permeation mechanism for the electron transfer at these hydroxy thiol derivatized electrodes should be addressed. In this mechanism, the redox centers would partition within the monolayer film, allowing the electron transfer to proceed across a smaller distance than entire monolayer film. Because the permeation of ions would be driven by the electric field present within monolayer, one would observe that at a given electrode potential the permeation rate (and hence the rate of electron transfer) would decrease with increasing thickness of the monolayer due to the decreasing electric field present within the monolayer. The monolayer could thus be free from pinholes or other defects without requiring electron tunneling through the entire monolayer thickness. While the dependence of the electron transfer rate with increasing thickness of the monolayer film shown in Figure 10 could be used to argue that permeation through these monolayer is occurring, the potential dependence of such a permeation would have to be quite contrary to simple expectations. In Figure 6, one sees that the reduction of Fe(cN)6+ is more facile than the reduction of Fe3+. As one polarizes the electrode negatively, the electric field within the monolayer will become more negative!' One would then anticipate that the cation, Fe3+,would (40) Charlot. G. Les Riactions Chimiques en Solution; Masson et Cie: Paris, 1969; Chapter 3, p 238.
Miller et al. permeate the monolayer better than the anion, Fe(CN)63-. A similar observation can be made for the shorter hydroxy thiols comparing the Tafel plots shown in Figure 6. If one neglects the sign of the electric field, the potential dependence of the observed electron-transfer rate would still not be consistent for the permeation reaction. At large electrode overpotentials, we see curvature in the Tafel plots. The rate of electron transfer becomes less strongly dependent on the electrode potential at high overpotentials. Because the electric field will increase continuously within this voltage range, the rate of field assisted permeation would have to become less dependent on the electric field at high overpotentials if it is to be consistent with the observed data. Such a dependence, while expected from the Marcus theory for the electron transfer, would be not expect for the permeation of ions through the monolayer film. A final indication of the low level of defects in these hydroxy thiol layers is their stability to chemical attack by sulfochromic acid. It is a striking observation that a hydroxy thiol monolayer on Au can withstand several minutes exposure to room-temperature sulfochromic acid without any change in its electrochemical properties. Maoz and Sagiv have shown that double bonds in various monolayer preparations can show considerable resistance to oxidation by permanganate ions!** The presence of defects within the monolayer and the lateral mobility of the hydrocarbon chains were concluded to be required for the oxidation reaction. In our experience using alkanethiol monolayers (such as octadecyl mercaptan) self-assembled on Au electrodes, such a sulfochromic acid treatment would be sufficient to remove the monolayer film. As a further test of the monolayer stability, the HO(CH2),,SH electrode was exposed to hot sulfochromic acid. As seen in Figure 9, this exposure to the hot sulfochromic acid does not cause a dramatic improvement in the Fe(CN)d- reduction kinetics even though the electrode capacitance has increased significantly. Because the presence of only a small coverage of bare Au areas would result in the reversible voltammetry of the ferricyanide at the electrode, one can conclude that the sulfochromic acid exposure seems to be evenly oxidizing the hydroxyl terminal function, effectively shortening the length of the hydroxy thiol layer. While being roughly 20% thinner as measured by the capacitance increase (assuming the same relative dielectric constant), the monolayer remaining after the hot sulfochromic acid treatment seems to be equally effective in lowering the ferricyanide reduction kinetics as the initial monolayer. We interpret this as being due to the conversion of the terminal function of the monolayer from a hydroxyl to a carboxylic group. At the pH used in this experiment, such a carboxyl group would be negatively charged, increasing the Coulombic repulsion of the negative ferricyanide from the electrode surface. Acknowledgment. This research was supported by the Fondation Suisse pour la Recherche en Microtechnique. The authors also thank Dr.J. Cook and Mr. C. M. Lee for their help in the XPS and ellipsometric measurements, respectively. (41) Because we do not know the point of zero charge of these derivatized electrodes, we cannot specify the absolute field within the monolayer at any one potential. (42) Maoz, R.; Sagiv, J . Thin Solid Films 1985, 132, 135. (43) Maoz, R.; Sagiv, J. Langmuir 1987, 3, 1034. (44) Maoz, R.; Sagiv, J. Longmuir 1987, 3, 1045.
