508
Langmuir 1996, 12, 508-512
Influence of Electrolyte Activity on Formal Potentials Measured for Ferrocenylhexanethiol Monolayers on Gold: Indistinguishable Responses in Aqueous Solutions of HClO4 and NaClO4 Jody Redepenning* and Jacqueline M. Flood† Department of Chemistry, University of Nebraska, Lincoln, Nebraska 68588-0304 Received August 25, 1995X A detailed study of the influence of electrolyte activity on apparent formal potentials measured for self-assembled monolayers of ferrocenylhexanethiol at high coverages on gold electrodes is provided. Ion selective electrodes are used in a logical sequence of comparisons to eliminate liquid junction potentials and, ultimately, to provide a rigorous comparison of apparent formal potentials for the monolayer over a broad range of concentrations in aqueous solutions of HClO4 and NaClO4. The results obtained in HClO4 and NaClO4 solutions are indistinguishable. Under the conditions examined it appears that protons and sodium ions have little influence on the apparent formal potentials. Hence it seems that protons and sodium ions have little influence on the overall interfacial potential distribution and, thus, little influence on the overall ionic structure near the monolayer/solution interface. The picture that emerges for the oxidized “monolayer” is one in which the high surface charge produced by the high surface coverage of ferroceniums is compensated by a highly compact layer of counterions.
Introduction The propensity of functionalized alkanes to form organized assemblies such as Langmuir-Blodgett films and self-assembled layers of organosulfur compounds on gold has made it possible to tailor surface structures to exhibit a wide variety of properties.1-27 These systems find applications ranging from fundamental studies of electron transfer to tribology and corrosion prevention. * To whom correspondence should be addressed. † Present address: Cargill Inc., Blair, NE. X Abstract published in Advance ACS Abstracts, December 15, 1995. (1) Stern, D. A.; Wellner, E.; Salaita, G. N.; Laguren-Davidson, L.; Lu, F.; Batina, N.; Frank, D. G.; Zapien, D. C.; Walton, N.; Hubbard, A. T. J. Am. Chem. Soc. 1988, 110, 4885. (2) Sasaki, T.; Bae, I. T.; Scherson, D. A.; Bravo, B. G.; Soriaga, M. P. Langmuir 1990, 6, 1234. (3) (a) Chidsey, C. E. D.; Bertozzi, C. R.; Putvinski, T. M.; Mujsce, A. M. J. Am. Chem. Soc. 1990, 112, 4301. (b) Chidsey, C. E. D. Science 1991, 251, 919. (4) Creager, S. E.; Collard, D. M.; Fox, M. A. Langmuir 1990, 6, 1620. (5) (a) Kunitake, M.; Akiyoshi, K.; Kawatana, K.; Nakashima, N. J. Electroanal. Chem. 1990, 292, 277. (b) Kunitake, M.; Kawahara, S.; Nakashima, N.; Manabe, O. J. Electroanal. Chem. 1991, 309, 341. (c) Kunitake, M.; Nasu, K.; Manabe, O.; Nakashima, N. Bull. Chem. Soc. Jpn. 1994, 67, 375. (6) (a) De Long, H. C.; Buttry, D. A. Langmuir 1990, 6, 1319. (b) De Long, H. C.; Donohue, J. J.; Buttry, D. A. Langmuir 1991, 7, 2196. (c) De Long, H. C.; Buttry, D. A. Langmuir 1992, 8, 2491. (d) Schneider, T. W.; Buttry, D. A. J. Am. Chem. Soc. 1993, 115, 12391. (7) Ueyama, S.; Isodo, S. J. Electroanal. Chem. 1991, 310, 281. (8) Lee, K. A. B. Langmuir 1990, 6, 709. (9) (a) Creager, S. E.; Rowe, G. K. Anal. Chim. Acta 1991, 246, 233. (b) Rowe, G. K.; Creager, S. E. Langmuir 1991, 7, 2307. (c) Creager, S. E.; Rowe, G. K. Langmuir 1993, 9, 2330. (d) Creager, S. E.; Weber, K. Langmuir 1993, 9, 844. (e) Rowe, G. K.; Creager, S. E. J. Phys. Chem. 1994, 98, 5500. (f) Creager, S. E.; Rowe, G. K. J. Electroanal. Chem. 1994, 370, 203. (10) Kwan, W. S. V.; Atanasoska, L.; Miller, L. L. Langmuir 1991, 7, 1419. (11) Acevedo, D.; Abruna, H. D. J. Phys. Chem. 1991, 95, 9590. (12) (a) Shimazu, K.; Ichizo, Y.; Sato, Y.; Uosaki, K. Langmuir 1992, 8, 1385. (b) Uosaki, K.; Sato, Y.; Kita, H. Langmuir 1991, 7, 1510. (13) Hickman, J. J.; Ofer, D.; Laibinis, P. E.; Whitesides, G. M.; Wrighton, M. S. Science 1991, 252, 688. (14) Obeng, Y. S.; Laing, M. E.; Friedli, A. C.; Yang, H. C.; Wang, D.; Thulstrup, E. W.; Bard, A. J.; Michl, J. J. Am. Chem. Soc. 1992, 114, 9943. (15) (a) Popenoe, D. D.; Deinhammer, R. S.; Porter, M. D. Langmuir 1992, 8, 2521. (b) Zak, J.; Yuan, H.; Ho, M.; Woo, L. K.; Porter, M. D. Langmuir 1993, 9, 2772. (16) Hickman, J. J.; Laibinis, P. E.; Auerbach, D. I.; Zao, C.; Gardner, T. J.; Whitesides, G. M.; Wrighton, M. S. Langmuir 1992, 8, 357.
0743-7463/96/2412-0508$12.00/0
With the goal of providing a more detailed structural picture of the monolayer/solution interface of charged selfassembled monolayers, we recently published a short manuscript that described the influence of supporting electrolyte activity on the apparent formal potentials measured for (pyridine)Ru(NH3)52+/3+ centers attached to gold electrodes at high surface coverage in self-assembled monolayers.21 For this system we proposed that a simple macroscopic electroneutrality argument might be used to account for shifts in apparent formal potentials. Because voltammetry of the self-assembled (pyridine)Ru(NH3)52+/3+ system is well-behaved in aqueous electrolytes in which chloride is the anion, it was possible to construct electrochemical cells in which there were no liquid junction potentials between the working and reference half cells. These factors allowed us to demonstrate that the apparent formal potentials measured for this system could be described in terms of a simple one-electron/one-counterion process. At high surface coverages the behavior is consistent with a special case of the model recently introduced by White and Smith28 concerning the influence of the interfacial potential distribution on the voltammetry of electroactive molecular films. (17) (a) Finklea, H. O.; Hanshew, D. D. J. Am. Chem. Soc. 1992, 114, 3173. (b) Finklea, H. O.; Rafenscroft, M. S.; Snider, D. A. Langmuir 1993, 9, 223. (c) Finklea, H. O.; Hanshew, D. D. J. Electroanal. Chem. 1993, 347, 327. (18) (a) Curtin, L. S.; Peck, S. R.; Tender, L. M.; Murray, R. W.; Rowe, G. K.; Creager, S. E. Anal. Chem. 1993, 65, 386. (b) Tender, L.; Carter, M. T.; Murray, R. W. Anal. Chem. 1994, 66, 3173. (c) Richardson, J. N.; Peck, S. R.; Curtin, L. S.; Tender, L. M.; Terrill, R. H.; Carter, M. T.; Murray, R. W.; Rowe, G. K.; Creager, S. E. J. Phys. Chem. 1995, 99, 766. (19) Katz, E.; Itzhak, N.; Willner, I. Langmuir 1993, 9, 1392. (20) Song, S.; Clark, R. A.; Bowden, E. F.; Tarlov, M. J. J. Phys. Chem. 1993, 97, 6564. (21) Redepenning, J.; Tunison, H. M.; Finklea, H. O. Langmuir 1993, 9, 1404. (22) Zhang, L.; Lu, T.; Gokel, G. W.; Kaifer, A. E. Langmuir 1993, 9, 786. (23) Fujihara, M. Photoinduced Electron Transfer and Energy Transfer in Langmuir-Blodgett Films. In Molecular and Biomolecular Electronics; American Chemical Society: Washington, DC, 1994; Vol. 240; pp 373-394. (24) Herr, B. R.; Mirkin, C. A. J. Am. Chem. Soc. 1994, 116, 1157. (25) Willicut, R. J.; McCarley, R. L. Langmuir 1995, 11, 296. (26) Sayre, C. N.; Collard, D. M. Langmuir 1995, 11, 302. (27) Everett, W. R.; Welch, T. L.; Reed, L.; Fritsch-Faules, I. Anal. Chem. 1995, 67, 292. (28) Smith, C. P.; White, H. S. Anal. Chem. 1992, 64, 2398.
© 1996 American Chemical Society
Ferrocenylhexanethiol Monolayers on Gold
In this article we describe the voltammetric response of ferrocenylalkanethiols self-assembled at high coverages on gold electrodes. This system was chosen for several reasons. Besides showing how electrolyte activity might be used to control the apparent formal potential observed for this particular monolayer, we sought a means of making unambiguous comparisons between voltammetric measurements made in different electrolyte solutions, the idea being that such measurements might ultimately be useful in obtaining a better understanding of interfacial structures and in interpreting the merits of descriptions of local solvation9e and interfacial potential distributions9e,28 that have recently appeared for self-assembled systems. To accomplish this we desired a self-assembling redox couple that is easily prepared, exhibits reversible voltammetry, and is reasonably stable in both oxidation states. Previous work by Creager et al.9 has demonstrated that ferrocenylalkanethiols meet these criteria, but because the voltammetry of ferrocenylalkanethiols systems is not well-behaved when chlorides are used as the electrolyte, it is difficult to measure rigorously the influence of electrolyte activity on the apparent formal potential. Liquid junction potentials between the working and reference electrodes change as the concentration of electrolyte changes; thus, interpretation of cell potentials is difficult. We have circumvented these problems in a manner similar to that described by Zielen.29 Cell potentials have been measured as a function of mean ionic activity and expressed mathematically using combinations of single ion activities: never measuring or knowing the single ion activities themselves. While it is true that single ion activities cannot be defined rigorously,30,31 it has also been noted that “...the concept may not the less be, and often has been, quite useful,32,33 when estimating activity coefficients in cases where great accuracy is not claimed.”30 This accuracy is certainly sufficient for the voltammetric measurements described below. As a result, by making a logical sequence of comparisons, we have been able to use ion selective electrodes to eliminate liquid junction potentials and to compare formal potentials for selfassembled ferrocenylalkanethiol monolayers in electrolytes containing different concentrations of different perchlorate salts. Experimental Section Equipment. All electrochemical measurements were performed in a water-jacketed isothermal cell maintained at 25 °C with a Neslab RTE-221 circulating bath. All of the cyclic voltammetry was performed with a Cypress Systems Model CYSY-1R potentiostat. An Orion Model 501 pH meter was used to measure cell potentials at zero current. Materials. 6-Ferrocenylhexanethiol was prepared according to a literature procedure.9f The surfaces of polycrystalline gold wires were prepared as follows. After a 10 min exposure to hot, concentrated nitric acid, electrodes were rinsed with distilled deionized water, air dried and then heated to incandescence in a gas/air flame. Finally, self-assembled films were prepared by exposing cleaned electrodes to 1 mM solutions of 6-ferrocenylhexanethiol in ethanol for at least 24 h. Perchloric acid concentrations were determined by conventional acid-base titration with a standard solution of sodium hydroxide. The sodium hydroxide solution was standardized against a potassium hydrogen phthalate primary standard. At electrolyte concentrations greater than 0.1 m, activity coefficients were obtained by interpolation of data tabulated by Stokes and Robinson34 and by Harned and Owen.35 At concentrations less than 0.1 m, activity (29) Zielen, A. J. J. Am. Chem. Soc. 1963, 67, 1474. (30) Kielland, J. J. Am. Chem. Soc. 1937, 59, 1675. (31) Guggenheim, E. A. J. Phys. Chem. 1929, 33, 842. (32) Guggenheim, E. A.; Schindler, T. D. J. Phys. Chem. 1933, 38, 533. (33) Scatchard, G. Chem. Rev. 1936, 19, 323. (34) Robinson, R. A.; Stokes, R. H. Trans. Faraday Soc. 1949, 45, 612.
Langmuir, Vol. 12, No. 2, 1996 509
Figure 1. Apparent formal potentials for a self-assembled monolayer and potentials for ion selective electrodes measured versus SSCE (a) Representative data collected in aqueous NaClO4 solutions: 0, SAM versus SSCE; O, Na+-ISE versus SSCE. (b) Representative data collected in aqueous HClO4 solutions: ], SAM versus SSCE; O, pH electrode versus SSCE. coefficients were calculated using the extended Debye-Huckel equation and ionic size parameters provided by Kielland.30
Results and Discussion It would be attractive to be able to use an ion selective electrode (ISE) as the reference electrode when a conventional three-electrode potentiostat is used to perform cyclic voltammetry. For example, if an acid were used as the supporting electrolyte, then a pH electrode would make a useful reference electrode. In a straightforward manner the influence of pH on the potential of the reference electrode could be subtracted from the overall cell potential to give the influence of electrolyte activity on the working electrode. However, practical difficulties make cyclic voltammetry essentially impossible when attempting to control the potential of a working electrode versus an ion selective glass electrode. The resistance across the glass membrane is prohibitively large. Consequently, an indirect method is needed to determine the potential of a working electrode versus a glass reference. This can be accomplished by using a reference electrode such as a saturated calomel electrode (SCE) to run a voltammogram in the usual manner. Upon completion of the voltammogram the potentiostat can be disconnected and, in the same solution, the potential can be measured between a desired ion selective electrode and the SCE reference. Both the electrolytic and the galvanic cells contain a liquid junction, and the potentials that develop across this liquid junction change as the activity of the electrolyte is changed, but for any given solution the liquid junction potentials (35) Harned, H. S.; Owen, B. B. Appendix A. In The Physical Chemistry of Electrolytic Solutions; Reinhold Book Corporation: New York, 1958.
510
Langmuir, Vol. 12, No. 2, 1996
Redepenning and Flood
includes the asymmetry potential of the glass membrane and a term including the single ion activity of the cation in the internal filling solution. The value of Kref is an experimentally determined quantity which will be discussed later. Equation 1 can be rewritten as shown in eq 2
Eref ) Kref +
RT ln γ+m+ F
(2)
where γ+ and m+ are the single ion activity coefficient and the molality of the supporting electrolyte cation in solution. If the half reaction for the self-assembled ferrocenylhexanethiol monolayer takes the following form
(Fc+‚‚‚X-)surf + e- ) (Fc0)surf + X-soln Figure 2. Apparent formal potentials versus potential of ion selective electrodes. (a) SAM versus Na+-ISE in NaClO4 solutions. (b) SAM versus pH electrode in HClO4 solutions.
for the electrolytic and galvanic cells are identical. Thus, by subtracting the potential of the ion selective electrode (versus SCE) from that of the working electrode (versus the same SCE in the same electrolyte solution), the potential of the working electrode can be obtained versus the ion selective electrode in the absence of liquid junction potentials. Experimental results for representative gold electrodes modified with self-assembled monolayers of ferrocenylhexanethiol are shown in Figure 1a and b. Figure 1a shows the influence of the sodium perchlorate activity on the apparent formal potential, (E°′SAM)app, which is approximated by (Ep,a + Ep,c)/2, measured for a self-assembled ferrocenylhexanethiol monolayer versus a sodium chloride saturated calomel electrode (SSCE). The potential of a sodium ion selective electrode (Na+-ISE) measured versus SSCE is also shown for the same solutions. Figure 1b shows the influence of the perchloric acid activity on the apparent formal potentials measured for a self-assembled ferrocenylhexanethiol monolayer versus SSCE. The potential of a pH electrode measured versus SSCE in the same solutions is also shown. Figure 2 shows the influence of electrolyte activity on (E°′SAM)app for ferrocenylhexanethiol monolayers in sodium perchlorate and perchloric acid solutions versus Na+-ISE and pH electrodes, respectively. Apparent formal potentials for 13 different monolayers were determined versus a sodium ion selective electrode at a nominal sodium perchlorate concentration of 0.1 m. The standard deviation obtained for the average value at this concentration, 623.0 ( 5.3 mV, is indicative of the reproducibility with which the self-assembled monolayers were prepared. Of the 13 electrodes, 6 were examined over sodium perchlorate concentrations ranging nominally from 0.005 to 1 m. Pooled values for (E°′SAM)app - ENaISE are plotted versus log a(NaClO4) in plot a of Figure 2. Six different monolayers were prepared and examined in perchloric acid solutions with concentrations ranging nominally from 0.001 to 1 m. Pooled values for (E°′SAM)app - EpH are plotted versus log a(HClO4) in plot b of Figure 2. The slopes for both plots in Figure 2 are very close to -118.3 mV/decade, a value which strongly suggests that a simple set of half-reactions might be used to provide a relatively straightforward explanation for the response. Such an explanation is found below. The potential of a generic cation selective electrode can be written as shown in eq 1
Eref ) Kref +
RT ln a+ F
(1)
where a+ is the single ion activity of the electrolyte cation in the external solution and Kref is a constant which
(3)
then, provided that the gold electrode is at equilibrium with the monolayer, the potential of the working electrode is given by eq 4
ESAM ) E°SAM +
aFcX(surf) RT ln F aFc(surf)aX-(soln)
(4)
where aFcX(surf) is the surface activity of the surface-bound ferrocenium with its counterion, aFc(surf) is the surface activity of the surface-bound ferrocene, and aX-(soln) is the single ion activity of the anion in solution. Equation 4 can be written as shown in eq 5
ESAM ) E°SAM +
RT aFcX(surf) RT ln ln γ-m- (5) F aFc(surf) F
where γ- and m- are the single ion activity coefficient and the molality of the supporting electrolyte anion in solution. Since the mean ionic activity coefficient for a one-to-one electrolyte is defined as follows,
γ( ) (γ+γ-)1/2
(6)
eq 4 can be written as shown in eq 7 2
ESAM ) E°SAM +
RT aFcX(surf) RT γ( m ln ln F aFc(surf) F γ+ -
(7)
Because the concentrations of cations and anions are equal in one-to-one electrolyte solutions, m can be defined as
m ) m+ ) m-
(8)
Finally, eqs 2, 7, and 8 can be combined to give the cell potential
Ecell ) ESAM - Eref )
[
E°SAM - Kref +
]
RT aFcX(surf) 2RT ln ln γ(m (9) F aFc(surf) F
Equation 9 shows that if Ecell can be measured at a fixed {aFcX(surf)/aFc(surf)} in a number of electrolyte solutions at different concentrations, a plot of Ecell versus log γ(m should give a straight line with a slope of -118.3 mV at 25 °C. Because we do not know how nonidealities are distributed between γ+ and γ-, we cannot know that either of the two half-cell potentials will give a slope of -59.16 mV when their potentials are plotted versus log γ(m. For whole cells we do not need to know how nonidealities are distributed between γ+ and γ-, however. Provided that the proposed half-reaction for the monolayer is actually appropriate and that the glass electrode gives a Nernstian response, it must be true that Ecell versus log γ(m will give a slope of -118.3 mV regardless of how the nonidealities are distributed. But can we know that Ecell is being measured at a constant value of {aFcX(surf)/aFc(surf)} in each electrolyte solution? We certainly do not know this a priori. We cannot even establish that this is true “after the fact.”
Ferrocenylhexanethiol Monolayers on Gold
Langmuir, Vol. 12, No. 2, 1996 511
The formalism chosen gives no assurance that the surface activities remain constant as the electrolyte activity changes in solution. However, the results in Figure 2 do strongly suggest that the half-reaction in eq 3 accurately depicts ion transfer that is coupled to electron transfer between the electrode surface and the self-assembled monolayer. Any changes in surfaces activities, which might occur upon changing the electrolyte activity in solution, are either fortuitously self-compensating or small. If it is true that (E°′SAM)app depends only on the activity of perchlorate in solution, i.e., that it is independent of the activity of the electrolyte cation, then it should be possible to make a direct comparison between the values of (E°′SAM)app measured in HClO4 solutions with those measured in NaClO4 solutions. The responses should be indistinguishable. Under these circumstances differences in Kref (see eq 2) for the Na+ and H+ selective electrodes should account for the offset observed in the cell potentials plotted in Figure 2. A qualitative description of how we opted to compare half-cell potentials measured for ferrocenylalkanethiol monolayers in different electrolytes at different concentrations is provided below. The qualitative description is followed by a detailed algebraic argument. After determining (E°′SAM)app versus a Na+-ISE in NaClO4 solutions and versus a pH electrode in HClO4 solutions, potentials between the Na+-ISE and a Ag/AgCl electrode were measured in a series of NaCl solutions to determine the response of the Na+ selective electrode to sodium ion activity. The potential between the H+ selective electrode and a Ag/AgCl electrode was measured in a series of HCl solutions to determine the response of the pH electrode to proton activity. Because the potential of the Ag/AgCl electrode versus the hydrogen electrode has a known dependence on hydrochloric acid activity, the potential of each of the ion selective electrodes can be obtained versus the standard hydrogen electrode. In summary, knowing the potential of the modified gold electrode versus the Na+ and H+ selective electrodes, and knowing the potential of these ion selective electrodes versus the hydrogen electrode, makes it possible to determine (E°′SAM)app for the modified gold electrode versus the standard hydrogen electrode. In this manner the ISE half-cells can be mathematically eliminated to give cell potentials free of the uncertainties introduced by changing liquid junction potentials, and unambiguous comparisons can be made between cell potentials measured in different electrolyte solutions. A more rigorous analysis of the measurements made in perchloric acid solutions (Figure 2b) is obtained by expressing the potentials of appropriate cells as shown in eqs 10-12.
[
]
aFcClO4(surf)
RT ln F aFc0(surf) RT RT ln γH+mH+ ln γClO4-mClO4- (10) F F
ESAM - EpH ) E°SAM - KpH +
EpH - EAg/AgCl ) KpH - E°Ag/AgCl +
RT RT ln γH+mH+ + ln γCl-mClF F (11)
EAg/AgCl - EH2 ) E°Ag/AgCl - E°H2 -
RT RT ln γH+mH+ ln γCl-mClF F (12)
Equation 10 describes the potential of the self-assembled monolayer versus the pH electrode in aqueous solutions
of perchloric acid. It is a more specific form of the generic expression given in eq 9. Equation 11 gives the potential of the pH electrode versus a AgCl-coated Ag wire exposed directly to solutions of HCl. Equation 12 gives the potential of the Ag/AgCl half-cell versus a hydrogen electrode in the same HCl solutions. If all molalities in equations 10-12 are equal, then these three equations can be added to give the potential of the self-assembled monolayer versus the hydrogen electrode as shown in eq 13
[
]
RT aFcClO4(surf) ln F aFc(surf) RT RT ln γH+mH+ ln γClO4-mClO4- (13) F F
ESAM - EH2 ) E°SAM - E°H2 +
or, given eq 6,
[
]
RT aFcClO4(surf) ln F aFc(surf) (0.1183) log(aHClO4) (14)
ESAM - EH2 ) E°SAM - E°H2 +
A similar set of equations can be used to interpret the measurements originally made in sodium perchlorate solutions (Figure 2a).
ESAM - ENaISE )
[
]
RT aFcClO4(surf) ln F aFc(surf) RT RT ln γNa+mNa+ ln γClO4-mClO4- (15) F F
E°SAM - KNaISE +
ENaISE - EAg/AgCl ) KNaISE - E°Ag/AgCl +
RT RT ln γNa+mNa+ + ln γCl-mClF F (16)
Equation 15 describes the potential of the self-assembled monolayer versus the sodium ion selective electrode in aqueous solutions of sodium perchlorate. It is also a more specific form of the generic expression given in eq 9. Equation 16 gives the potential of the sodium ion selective electrode versus a AgCl-coated Ag wire exposed directly to solutions of NaCl. Equation 12 is again needed to give the potential of the Ag/AgCl half-cell versus a hydrogen electrode in solutions of HCl. If all molalities in eqs 15, 16, and 12 are equal, then these three equations can be added to give the potential of the self-assembled monolayer versus the hydrogen electrode as shown in eq 17
[
]
RT aFcClO4(surf) ln F aFc(surf) (0.1183) log(aHClO4) (17)
ESAM - EH2 ) E°SAM - E°H2 +
which is identical to eq 14. Equations 10-17 give direction concerning the interpretation of experimental data shown in Figure 2. First, it is necessary to measure potentials for the Na+-ISE and pH electrodes over a range of NaCl and HCl concentrations, respectively, versus a Ag/AgCl reference that is exposed directly to the same electrolyte solutions. The results of such measurements are plotted in Figure 3 versus log a, where a is the activity of the appropriate electrolyte. The slopes of the plots in Figure 3 immediately suggest an explanation for the small but significant difference observed from the slopes in Figure 2. The pH electrode we used exhibited a “super-Nernstian” response to hydrogen ions. While this response is somewhat of a distraction, it does not prevent us from providing a rigorous comparison between values of (E°′SAM)app measured in HClO4 and NaClO4 solutions.
512
Langmuir, Vol. 12, No. 2, 1996
Figure 3. Potentials of ion selective electrodes measured versus a Ag/AgCl electrode exposed directly to electrolytes containing chloride. (a) Na+-ISE versus Ag/AgCl in NaCl solutions. (b) pH electrode versus Ag/AgCl in HCl solutions.
The means by which eq 12 can be used in conjunction with the experimental data in Figures 2 and 3 requires further explanation. Consider any one of the data points in Figure 2a. (E°′SAM)app versus EH2 is not obtained simply by adding (E°′SAM)app versus ENaISE at a certain activity in Figure 2, ENaISE versus EAg/AgCl (from a fit to the data Figure 3), and EAg/AgCl versus EH2 at the same activity. The cell potentials must be compared at the same molality as that of the original NaClO4 solution, not at the same mean ionic activity as the original NaClO4 solution. For example, one of our ferrocenylhexanethiol monolayers gave a value of +0.273 V for (E°′SAM))app - ENaISE when the NaClO4 concentration was 0.0992 m (when a(NaClO4) was 0.0775 m). When the concentration of NaCl is 0.0992 m, the activity of NaCl is 0.0772 m, and ENaISE - EAg/AgCl is calculated to be -0.0004 V from the fit to the data shown in Figure 3a. When the concentration of HCl is 0.0992 m, the activity of HCl is 0.0791 m, and EAg/AgCl - EH2 is calculated to be 0.3526 V using eq 12. Adding these values of (E°′SAM)app - ENaISE, ENaISE - EAg/AgCl, and EAg/AgCl - EH2 gives a value of 0.625 V for (E°′SAM)app - EH2. Repeating this process for each data point collected in the various HClO4 and NaClO4 solutions allows us to compare the experimental data with the result predicted by eqs 14 and 17. When the same rationale used to derive identical eqs 14 and 17 is applied to the raw experimental data originally collected in HClO4 and NaClO4 solutions, it is not possible to distinguish between values of (E°′SAM)app measured in the two electrolytes. A linear least squares fit to the data collected in HClO4 solutions is within one standard deviation of the fit to the data collected in NaClO4 solutions. A fit to (E°′SAM)app - EH2 versus log a(HClO4) for data originally collected in HClO4 gives a slope of -116.1 ( 1.3 mV and an intercept of 489.7 ( 2.3 mV. A fit to (E°′SAM)app - EH2 versus log a(HClO4) for data originally collected in NaClO4 gives a slope of -116.5 ( 1.4 mV and an intercept of 492.6 ( 2.0 mV. In both cases the intercepts are the formal potential for the selfassembled monolayer, E°′SAM, measured versus the standard hydrogen electrode. Because there does not appear to be a significant difference between the two data sets, all of measurements of (E°′SAM)app in both electrolytes were pooled and plotted as shown in Figure 4. The appearance that many more measurements were made in HClO4 solutions than in
Redepenning and Flood
Figure 4. Apparent formal potentials for self-assembled monolayers versus the hydrogen electrode: 0, data originally collected in NaClO4; ], data originally collected in HClO4.
NaClO4 solutions is an artifact caused by plotting the data collected in solutions of HClO4 “on top” of that collected in solutions of NaClO4. While the slope of the plot in Figure 4 closely approaches the ideal slope of -118.3 mV, the difference between the observed slope and the ideal slope is significant. The reason for this small difference is not clear. One alternative is that the simple oneelectron/one-counterion half-reaction shown in eq 3 accurately depicts the actual electrochemical process, but that the surface activities change to some small extent as the electrolyte activity changes in solution. Another alternative is that the one-electron/one-counterion model is not completely accurate because the self-assembled monolayer is not completely charge compensated by perchlorate ions. This alternative is essentially that which is treated in the description recently provided by Smith and White.28 While the ideal slope predicted in eqs 14 and 17 appears to be significantly different from that actually observed, it is important not to dwell extensively on the small difference. It is important to note that over 3 orders of magnitude the difference between the ideal slope and the observed slope is only ∼2.5 mV, a value that compares favorably with popular ion selective electrodes. The selfassembled ferrocenylhexanethiol monolayers examined here certainly have limited ability to discriminate between different types of anions, but under the conditions we have examined, it appears that nearly every ferrocenium is charge compensated by a perchlorate ion in close proximity. Protons and sodium ions appear to have little influence on the apparent formal potentials observed for the monolayer. Hence it seems that these ions have little influence on the overall interfacial potential distribution and, thus, little influence on the overall ionic structure near the interface. The picture that emerges for the oxidized “monolayer” is one in which the high surface charge produced by the high surface coverage of ferroceniums is compensated by a highly compact layer of counterions. The influence of counterion activity on electrode potential takes a form which is virtually identical to the form for Ag/AgCl and Hg/Hg2Cl2 half-reactions. Acknowledgment. J.R. is pleased to acknowledge NSF EPSCoR Cooperative Agreement OSR-9255225 for support of this research. LA950715T
