2307
Langmuir 1991, 7, 2307-2312
Redox and Ion-Pairing Thermodynamics in Self -Assembled Monolayers Gary K. Rowe and Stephen E. Creager' Department of Chemistry, Indiana University, Bloomington, Indiana 47405 Received May 9, 1991. In Final Form: June 6, 1991 Self-assembled mixed monolayers composed of ferrocenylhexanethiol ((Fc(CHz)&H, FC = (q6-4H5)Fe($-CsH,)) and n-alkanethiolshave been prepared on polycrystalline bulk gold electrodes. Monolayers containing ferrocenylhexanethioland n-alkanethiolsfrom 4 to 12 carbons long show the following trends as n-alkanethiolchain length increases: (a) the quantity of immobilized ferrocene decreases; (b) the redox potential for ferrocene oxidation shifts to more positive values. The positive shift is explained by the progressively more alkane-like environment created by the coadsorbed alkanethiols, which destablizes ferricinium relative to ferrocene. The redox potential was also found to depend on electrolyte, shifting negative with log [HClOd]. Effective formation constants for forming ion pairs between ferricinium and several anions are calculated and compared; the values scale roughly with the energies of solvation of the anions. acids on metal oxide surfaces,28phosphonates on phosphonate surfaces,m and isocyanides on platinum.g0 ExMonolayer self-assembly chemistry is a most promising posure of a gold surface to a dilute (1.0 mM) solution of strategy for constructing stable, well-defined monolayers an n-alkanethiol results in a chemisorbed monolayer that on electrode surfaces. The self-assembly technique works is densely packed in two dimensions and excludes ions well for molecules that bond strongly to a surface and that and water from the underlying g ~ l d . ~ JThe ~ Jthermo~ have shapes that pack well in two dimensions. Chemical dynamically favorable formation of a gold-thiolate bond strategies for self-assembly include thiols, disulfides, and makes the gold-thiol system ideal for monolayer selfsulfides on g0ld,'-~5silanes on silicon d i o ~ i d e , 2fatty ~ ~ ~ ~ assembly schemes, and the stability of that bond over a wide range of applied potential makes such a system * Author to whom correspondence should be addressed. suitable for electrochemical studies. Self-assembly chem(1) Chidsey, C. E. D.; Bertozzi, C. R.; Putvinski, T. M.; Mujsce, A. M. istry offers significant advantages over other approaches J. Am. Chem. SOC.1990,112, 4301. to electrode modification involving polymer films, which (2) Creager, S. E.; Rowe, G. K. Anal. Chim. Acta 1991,246, 233. (3) Chidsey, C. E. D. Science 1991,251,919. are usually much thicker and have considerable tertiary (4) Collard, D. M.; Fox, M. A. Langmuir 1991, 7,1192. structure, or transferred Langmuir-Blodgett films, which (5) Buttry, D. Langmuir 1990,6, 1319. often contain many defects and can be intrinsically (6) Creager, S. E.; Collard, D. M.; Fox, M. A. Langmuir 1990,6,1617. (7) Soriaga, M. P.; Bravo, B. J.; Michelhaugh, S. L. J. Electroanal. unstable. Chem. Interfacial Electrochem. 1988,241, 199. Self-assembled monolayers containing immobilized re(8) Obeng, Y. S.; Bard, A. J. Langmuir 1991, 7, 195. dox-active molecules provide an opportunity to study re(9) Porter, M. D.; Bright, T. B.; Allara, D. L.; Chidsey, C. E. D. J.Am. Chem. SOC.1987,109,3559. dox processes within a well-defined chemical environ(10) Trevor, D. J.; Chidsey, C. E. D.; Loiacono, D. N. Phys. Reu. Lett. ment.1-s Kinetics and thermodynamics of interfacial re1989,62 (8), 9295. dox reactions are known to be strongly affected by the (11) Wiechers, J.; Twomey, T.; Kolb, D. M. J. Electroanal. Chem. Interfacial Electrochem. 1988, 248,451. nature of the medium in which they occur. A promising (12) (a) Chidsey, C. E. D.; Loiancono,D. N.Langmuir 1990,6,682. (b) strategy for studying local medium effects on electrode Chidsey, C. E. D.; Liu, G.-Y.; Rowntree, P.; Scoles, G. J. Chem. Phys. reactions is to form monolayers in which a redox-active 1989, 91, 4421. (c) Chidsey, C. E. D.; Lui, G.-Y.; Scoles, G.;Wang, J. Langmuir 1990, 6, 1804. moiety is included as the minority component in a two(13) Creager, S. E., Indiana University, unpublished results. component self-assembled monolayer film. Structure and (14) (a) Finklea, H. 0.; Avery, S.; Lynch, M.; Furtach, T. Langmuir composition of other mixed monolayers formed from 1987,3,409. (b) Finklea, H. 0.; Snider, D. A.; Fedyk, J. Langmuir 1990, solutions of two different electroinactive, terminally 6, 371. (15) (a) Sabatani, E.; Rubinstein, I.; Maoz, R.; Sagiv, J. J.Electroanal. substituted alkanethiols have been shown to correlate with Introduction
Chem. Interfacial Electrochem. 1987,219, 365. (b) Sabatani, E.; Rubinstein, I. J. Phys. Chem. 1987,91,6663. (c) Rubinstein, I.; Steinberg, S.; Tor, Y.; Shanzer, A.; Sagiv, J. Nature 1988, 332, 426. (16) Fabianowski, W.; Coyle,L. C.; Weber, B.A.; Granata, R. D.;Castner, D. G.;Sadownik, A,; Regen, S. L. Langmuir 1989,5, 35. (17) Bunding Lee, K. A. Langmuir 1990,6,709. (18) Barner, B. J.; Corn, R. M. Langmuir 1990,6, 1023. (19) (a) Nuzzo, R. G.;Fueco, F. A.; Allara, D. L. J. Am. Chem. SOC. 1987,109,2358. (b) Nuzzo, R. G.; Dubois, L. H.; Allara, D. L. J. Am. Chem. SOC.1990,112,558. (c) Dubois, L. H.; Zegarski, B. R.; Nuzzo, R. G. J. Am. Chem. SOC.1990,112,570. (d) Dubois, L. H.; Zegarski, B. R.; Nuuo, R. G. Proc. Natl. Acad. Sei. U.S.A. 1987,84,4739. (e) Nuzzo, R. G.; Zagarski, B. R.; Dubois, L. H. J. Am. Chem. SOC.1987,109,733. (20) (a) Bain, C. D.;Troughton,E. B.; Tao,Y. T.; Evall, J.; Whitesides, G. M.; Nuzzo, R. G.J. Am. Chem. SOC.1989,111,321. (b) Whitesides, G. M.; Laibinis, P. E. Langmuir 1990,6, 87. (c) Bain, C. D.; Evall, J.; Whitesides, G. M. J. Am. Chem. SOC.1989, 111, 7155. (d) Bain, C. D.; Whitesides, G.M. J. Am. Chem. SOC.1989, 111, 7164. (e) Bain, C. D.; Whitesides, G.M. Angew. Chem., Intl. Ed. Engl. 1989,28,506. (21) Bain, C. D.; Biebuyck, H. A.; Whitesides, G.M. Langmuir 1989, 5 , 723. (22) Strong, L.; Whitesides, G. M. Langmuir 1988, 4, 546.
(23) Bain, C. D.; Whitesides, G.M. J.Am. Chem. SOC.1988,110,6660. (241 E. B.: Bain., C. .--,Trounhton. ---- D.: Whitesides. G. M.:' Nuzzo. R. G.: Allara, D. L.yPorter, M. D. Langmuir 1988,4,365. (25) Bain, C. D.; Whitesides, G. M. Science 1988,240, 62. (26) (a) Sagiv, J. J. Am. Chem. SOC.1980, 102, 92. (b) Netzer, L.; Sagiv, J. J. Am. Chem. SOC.1983, 105, 674. (c) Maoz, R.; Sagiv, J. J. Colloid Interface Sci. 1984.109.465. (dl Gun. J.: Iscovici. R.: W v , J. J. Colloid Interface Sei. 1984,101,201~(e) Gun, J.; Sagiv; J. J. d l l o i d Interface Sci. 1986, 112,467. (0Maoz, R.; Sagiv, J. Langmuir 1987,3, 1034. (9) Maoz, R.; Sagiv, J. Langmuir 1987,3, 1045. (27) (a) Tillman, N.; Ulman, A,; Penner, P. L. Langmuir 1989,6,101.
.- - . -
-
(b) Tillman, N.;Ulman, A.; Schildkraut, J. S.;Penner, T. L. J.Am. Chem.
SOC.1988,110,6136. (28) (a) Allara, D. L.; Nuzzo, R. G. Langmuir 1985,1,45. (b) AUara, D. L.; Nuzzo, R. G. Langmuir 1985, I, 52. (29) Lee, H.; Kepley, L. J.; Hong, H.-G.; Akhter, S.; Mallouk, T.E. J. Phys. Chem. 1988, 92, 2597.
(3O).Hickman,J. J.; Zou, C.; Ofer, D.; Harvey, P. D.; Wrighton, M. S.; Laibinis, P. E.; Ban, C. D.; Whitesides, G.M. J. Am. Chem. SOC.1989, 111, 7271.
0 1991 American Chemical Society
2308 Langmuir, Vol. 7, No. 10, 1991
(a) the relative concentration of each component in the coating solution, (b) the relative chain length of the two components, (c) the nature of the coating solvent, and (d) the nature of the terminal group (size, polarity, bonding, etc.).+25 Similar effects are likely for mixed monolayers formed from electroactive molecules. Redox-active molecules that have been prepared and studied in selfassembled monolayers include ferrocene,I+ viologen,5i6 hydroquinone,’ and Run(bpy)3.8 Ferrocene is an extremely well-characterized, neutral, reversible outer-sphere redoxactive molecule and is therefore an ideal choice for incorporation into self-assembled mixed monolayers. Further, the chemical modification of ferrocene via covalent attachment of various molecular linkages containing a thiol group is straightforward and has recently been demonstrated.l+ In this paper, we describe a study on self-assembled mixed monolayers containing ferrocenylalkanethiols and n-alkanethiols immobilized on gold. These investigations involved incorporation of ferrocenylhexanethiol ((Fc(CH2)6SH, Fc = (15-C5H5)Fe(15-C5H4))into monolayers containing n-alkanethiols from 4 to 12 carbons long. The effects of mole fraction of each component in the coating solution and of the relative chain lengths of the coadsorbates have been investigated. Also, several electrolytes were investigated to determine any changes they induce in the electrochemical response. The critical findings are that (i) the adsorbate with the longer alkyl chain is always preferentially adsorbed and (ii) as the coadsorbed alkanethiol chains become longer than the ferrocenylalkanethiol chain, the ferrocene moiety experiences a progressively more alkane-like environment. This local medium effect is manifested in a positive shift in redox potential for ferrocene oxidation and an increase in the propensity of electrogenerated ferricinium to ion pair with certain electrolyte anions. Experimental Section Materials. Substrates were polycrystalline gold (JohnsonMatthey, 99.998576) cut into 9 X 9 X 1mm squares and prepared as described below. Water was purified via a Barnstead Nanopure system. All solvents and electrolytes were of commercial origin and were used as received except for KPF6, which was recrystallized from aqueous alkaline solution. The n-alkanethiols were obtained from Aldrich and used as received. The ferrocenylhexanethioladsorbate was synthesized as follows: first, ferrocene was acylated under standard Friedel-Crafts conditions (CHZClz, 1equiv with 6-bromohexanoylchlorideto produce the intermediate ferrocenyl5-bromopentyl ketone in 83% yield (based on ferr~cene)?~ The carbonyl moiety was fully reduced via Clemmenson reduction (reflux over zinc amalgam in a biphasic mixture of toluene and aqueous 6 M HCl) to give the product ferrocenyl-6-bromohexane in 89 % yield. Finally, the bromide was converted to a thiol by reaction with a stoichiometric quantity of thiourea in ethanol, followed by basic (aqueous KOH) cleavage of the intermediate isothiuronium salt.%,= The product ferrocenylhexanethiol was obtained in 83% yield after purification by flash chromatography (silica gel, 1OO:l hexanes/ ethyl acetate). Structures of all isolated intermediates and products were confirmed by electron impact mass spectrometry and by ‘HNMR spectroscopy, and the final ferrocenylhexanethiol product was confirmed by microanalysis. Anal. Calcd for CleHztSFe: C, 63.58; H, 7.34; S, 10.61. Found: C, 63.26; H, 7.48; s, 10.99. (31) Hupp, J. T.Inorg. Chem. 1990,29, 5010. (32) Nelson, 1. V.; Iwamoto, R. T.Anal. Chem. 1963,35,867. (33) Vogel, M.;Rausch, M.; Rosenberg, H. J. Org. Chem. 1957, 22, 1016. (34) Cossar, B.C.;Fournier, J. 0.;Fields, D. L.;Reynolds,D. D.J . Org. Chem. 1962,27,93. (35) Wardel, J. L.In The Chemistry of the Thiol Group; Patai, S., Eds.;Wiley: London, 1974; Chapter 4.
Rowe and Creager Monolayer Formation. Monolayers are formed by soaking gold substrates in a coating solution containing ferrocenylhexanethiol and a nonsubstituted n-alkanethiol(l.0 mM total thiol concentration) in absolute ethanol. The samples are removed from the coating solutions after a period of no less than 11hand washed with copious amounts of ethanol followed by water. Prior to coating, the electrodes were potted in acrylic and polished with aqueous alumina slurries. After removal from the potting compound, they were skid polished by hand to a mirror finish with aqueous slurries of 1.0- and 0.3-pm alumina. They were further cleaned by etching for approximately 10 min in hot (100 “C)concentrated nitric acid in a quartz test tube and transferred (after rinsing with water and ethanol) directly into the coating solution. Electrochemical Measurements. The electrochemical cell is made from Kel-F machined to accept a clean Viton O-ring (9/32 in. diameter, 1/16 in. thick) in a cylindrical cavity surrounding a hole in the bottom of the cell and define an electrode area of approximately 0.4 cm2.2 The gold electrode is pressed against the O-ring, creating an effective seal and reproducibly defining the area exposed to the solution. Platinum wire auxiliary, Ag/ AgCl/saturated KCl reference, and Hg/Hg#O,/saturated KzSO4 reference electrodes were used exclusively. All solutions were purged with water-saturated argon (SealProducta Co.) prior to electrochemical data acquisition. Cyclic voltammetric measurements were made on a EG&G-PAR Model 362 scanning potentiostat. All experiments were done at room temperature.
Results and Discussion All experiments involved the following sequence of steps: (a) mechanical polishing and chemical etching of a gold electrode; (b) immersion in an ethanol solution containing ferrocenylhexylthiol and an n-alkanethiol for a minimum of 11 h; (c) extensive rinsing with absolute ethanol followed by distilled deionized water; (d) characterization by cyclic voltammetry in aqueous electrolyte solutions. Cyclic voltammetry was used to characterize these mixed monolayers since it allows for the determination of the quantity of immobilized ferrocene (r,in mol/ cm2by integration of the voltammetricpeak) and the redox potential ( E O ’ , taken as the average of the anodic and cathodic peak potentials). The quantity of immobilized ferrocene incorporated into these mixed monolayers is related to the mole fraction of each component in the coating solution. Figure 1shows how coverage varies with coating solution composition for a series of mixed monolayers of ferrocenylhexanethiol with butane-, hexane-, octane-, and decanethiol. These curves may be thought of as adsorption isotherms, although their rigorous thermodynamic basis has yet to be proved. The dotted line indicates the isotherm expected if interactions among adsorbates are minimal and adsorption of the electroactive and electroinactive coadsorbates is purely statistical; ferrocene coverage would then scale linearly with the mole fraction of ferrocene in the coating solution and, at a mole fraction of one, would produce a calculated maximum coverage of 4.5 X 10-10mol/cm2. This maximum coverage is obtained by assuming hexagonal close packing and by treating ferrocene as a sphere of diameter 6.6 A;% it is just slightly lower than our experimentally observed value of 4.6 X mol/cm2, confirming the presence of a single monolayer of adsorbed ferrocene under these conditions. Chidsey and co-workers reported a larger maximum ferrocene coverage of 5.8 X 10-lo mol/cm2 for a monolayer of slightly different structure (i.e. longer alkyl chain). We ascribe our better agreement with the calculated coverage to the fact that our molecule is shorter and therefore less apt to fold over and pack among the poorly packed polymethylene chains. Differences in surface roughness may also play a role. (36) Seiler,P.; Dunitz, J. D. Acta Crystallogr., Sect E 1979,E35,1068.
Langmuir, Vol. 7, No.10, 1991 2309
Self- Assembled Monolayers I
0
Butane
v Hexane
Octane o Decane v
I
1 I
Q)
0,
Y Q)
> 0
0
" W , J , 100
06
0.2
0.4
XFC
(in solution)
0.8
1.0
Figure 1. Ferrocene coverage vs mole fraction of ferrocene in ethanol coatin solution for mixed monolayers containing (Fc)CsSH and n-al%lanethiols:(A) n-butanethiol; (B)n-hexanethiol; (C) n-octanethiol; and (D)n-decanethiol. Coverage obtained by manual integration of cyclic vo1ta"ograms. Dashed line is ferrocene coverage expected for ideal statistical adsorption. Standard error in ferrocene coverage of selected data points are representative of all monolayer compositions shown. Inspection of Figure 1reveals that coadsorption offerrocenylhexanethiol and hexanethiol from ethanol is nearly ideal. For coadsorption of ferrocenylhexanethiol with alkanethiols shorter than six carbons, ferrocene adsorption is favored. In contrast, alkanethiol adsorption is favored for alkanethiols greater than six carbons in length. The tendency is for the adsorbate with the longest alkyl chain to always be preferentially adsorbed. This trend was noted by Chidsey and co-workers for similar mixed monolayers containing ferrocene-terminated alkanethiols and n-alkanethiols' and in a related electroinactive system by Bain and co-workers.2h It is tempting to ascribe these results to a favorable self-interaction parameter, perhaps relating to greater van der Waals interactions, for the longer-chain adsorbates. It is not yet clear however whether these monolayers form under kinetic or thermodynamic control and whether they can be fully described by normal adsorption isotherms. Purely thermodynamic adsorption should be dominated by the strength of the gold-thiolate bond and purely kinetic adsorption should correlate with the relative rates of diffusion of thiols in the coating solvent. To explain the adsorption process, we postulate a two-step model in which a "preformed" physisorbed layer forms initially, followed by formation of the gold-sulfur bond. Previous self-assemblystudies have used relative solubility and favorable interaction parameters arguments to explain the adsorption behavior of mixed monolayers;20cpdthese arguments are consistent with the proposed two-step model in that the composition of the preformed, physisorbed layer would be determined by these factors. The adsorption behavior a t very high coverages offerrocenylhexanethiol is complicated by the presence of several different adsorption domains as evidenced by multiple peaks in the voltammetry as shown in Figure 2. This may explain the irregularities in surface coverage of ferrocene a t these high coverages. We have performed most of our experiments at lower coverages to avoid these complications. Since the adsorption behavior with respect to relative adsorbate concentration is known for each pair of coadsorbates, it is possible to compare the electrochemical behavior in a series of monolayers, all with similar ferrocene coverage, in which the coadsorbate alkanethiol chain length is varied. The arrow a t the bottom right of
0.0
0.2
0.4
0.6
0.8
1.0
E (V vs. Ag/AgCI) Figure 2. Cyclic voltammogram in 1.0 M HClOd of an electrode coated from an ethanol solution containing (Fc)C&3Hand n-hexanethiol: scan rate, 100 mV/s; T,2.5 pA; coverage of ferrocene, 2.9 X mol/cm" (two-thirds of a monolayer).
Figure 1 shows the coverage we have chosen. This approach ensures that the voltammetry in each case represents reactivity of isolated ferrocene sites and removes possible artifacts caused by differing coverages (such as differing lateral interactions among redox-active sites). Figure 3 shows a series of cyclic voltammograms in 1.0 M HClOI taken for monolayers containing approximately 2.4 X 10-l' mol/cm2 (5% of a full monolayer) of ferrocenylhexanethiol coadsorbed with butane-, hexane-, octane-, decane-, or dodecanethiol. The electrode potential was scanned from -0.1 to 0.8 V and back at 100 mV/s. No loss of ferrocene adsorbate is observed upon repeated cycling in acidic aqueous medium for several minutes, demonstrating that the monolayers are stable to electrochemical cycling. The shape of each voltammogram is approximately independent of scan rate, and the peak current scales with scan rate as expected for an immobilized redox couple. There is less than 5-mV peak splitting in each voltammogram indicating that electron transfer is rapid on the time scale of these experiments. The anodic and cathodic peak currents are nearly identical for each voltammogram, confirming that electron transfer is chemically reversible. The most significant observation regarding Figure 3 is that, as the alkanethiol coadsorbate chain length increases, there is a corresponding shift in redox potential to more positive values. The shift is clearly thermodynamic in nature given that the anodic and cathodic peak potentials both shift together. A monolayer containing ferrocenylhexanethiol and coadsorbed dodecanethiol exhibits a redox potential almost 200 mV more positive than one containing coadsorbed hexanethiol. We propose that the shift in redox potential is due to a progressively more alkane-like environment around the ferrocene moiety; this environment would stabilize ferrocene to a greater extent than it would ferricinium; hence, we observe a shift in E O ' . This can be thought of as the thermodynamic "price" that must be paid to force the electrogenerated ferricinium ion to reside in this quite nonpolar, poorly solvating environment. Ferrocene oxidation has been observed previously to shift to more positive potentials where the electrochemical experiments were done with a series of n-alcoholsas The monolayer appears to be more permeable to electrolyte ions and/or solvent after oxidation since the capacitive charging current is always larger a t potentials more positive of the ferrocene redox potential. A second observation regarding Figure 3 is that the peak
Rowe and Creager
2310 Langmuir, Vol. 7,No.10, 1991
IS v
Hexane
‘HCIO,
Figure 5. Effect of perchlorate ion on the redox potential for electrodes coated from ethanol solutions containing (Fc)CeSH and n-alkanethiols: (A) n-butanethiol; (B)n-hexanethiol; (C) n-octanethiol;(D) n-decanethiol;(E)n-dodecanethiol. Scan rate = 100 mV/s. All solutions contain 1.0 M H2SO4; coverage of ferrocene is 2.4 X 10-11mol/cm2.
+ Figure 6. Model for ion pairing of surface-bound ferricinium with perchlorate.
0.7
0.6
-
0 0
05
0
0.1
L
2
4
6
8
10
I2
J
14
Number of Carbons Figure 4. Redox potentials of mixed monolayers containing (Fc)C&3Hand n-alkanethiols for voltammograms taken in 1.0M H2S04and 1.0 M HC104. Surface coverage of ferrocene is as in Figure 2. Scan rate = 100 mV/s.
shape becomes broad and asymmetric with increasing alkanethiol coadsorbate chain length, most likely due to a degree of disorder in the monolayer such that ferrocenes are located in a variety of slightly different environments. Figure 4 shows redox potential vs the number of carbons in the alkanethiol chain for monolayers containing ferrocenylhexanethiol in which the electrolyte was either 1.0 M HClO4 (triangles) or 1.0 M HzSO4 (circles). Inspection of Figure 4 reveals that the redox potential is always more positive in HzSO4 than in HClO4 electrolyte. Further, as the alkanethiol coadsorbate chain length increases beyond that of the ferrocene-containing species, the redoxpotential shift in HzSO4 is more dramatic than that in HC104 solution. This dependence of redox potential on the nature of the electrolyte is somewhat unexpected. We
find, for example, that the redox potential of the watersoluble analogue hydroxymethylferrocene is approximately +0.21 V (vs Ag/AgCl) independent of both the nature and concentration of electrolyte (slight shifts with changing electrolyte concentration are attributed to changing liquid junction potentials). This phenomenon, and the corresponding assumption that the ferrocene/ferricinium redox potential should be nearly invariant with solvent and electrolyte, has formed the basis of the socalled “ferrocene assumption” in which redox potentials of new compounds in unusual solvents can be profitably referenced against ferrocene/ferricinium rather than against an external reference electrode. Hence, we have investigated the dependenceof the immobilized ferrocene/ ferricinium redox potential on electrolyte ion further. Figure 5 is a plot of peak potential vs log [HClOd] (in the presence of 1.0 M HzSO4) for monolayers containing ferrocenylhexanethiolcoadsorbed with alkanethiols from 4 to 12 carbons long. The perchloric acid is added incrementally to a solution already containing a large concentration of sulfuric acid;this is necessary to maintain ionic strength, and therefore activity coefficientsand liquid junction potentials, as constant. At the far left of Figure 5, the x axis is broken to show the peak potentials in 1.0 M HzSO4 with no perchloric acid present. The peak potential shifts negatively for increasing amounts of perchloric acid added, resulting in a slope of approximately 54 mV per unit increase in log [HClOd] for each monolayer composition. We assume that the slope deviates slightly from the expected 59 mV because concentrations were used in place of activities. This behavior is consistent with consumption of a single perchlorate anion on ferrocene oxidation; we postulate that this represents formation of a discrete, specific ion pair between ferricinium and perchlorate (see Figure 6). Buttry suggested that C104- loses its solvation sphere to form specific ion pairs with viologen self-assembled mono-
Self-Assembled Monolayers
Langmuir, Vol. 7, No. 10, 1991 2311
Table I. Effective Formation Constants for Ion Pairing between Surface-ConfinedFerricinium and Perchlorate Anion. alkanethiol carbons 4 6 8
10 12 4.5
Eo’! V 0.34 0.35 0.44 0.55
0.62
rpc =
X
layers indicating the propensity of C104-to form ion pairs.5 Ion pairing of ferricinium could occur either from a “free” ferricinium ion or from a ferricinium that is already ion paired with bisulfate (the predominant anion in 1.0 M H2S04). In principle these two possibilites could be distinguished by the behavior with changing sulfuric acid concentration; our attempts along this line were unfortunately complicated by small changes in liquid junction potentials and/or activity coefficients as the electrolyte concentration was changed. The treatment below therefore considers data taken in the presence of a large concentration of as the background electrolyte. In the most general case, for simultaneous ion-pairing equilibria involving ferricinium with two different anions X and Y, it can be shown (see Appendix) that
Ep = E O ’ - (RT/nF)In (1 + KxCx + KyCy)
(1)
where E, is the peak potential (where half of the bound ferrocene has been oxidized) and KX and Ky are the formation constants for forming ion pairs of ferricinium with ions X and Y, respectively. (This treatment assumes no association of X or Y with neutral ferrocene.) The distribution of ferricinium among “free” Fc+, Fc+X-, and Fc+Y- is given by the relative magnitudes of the three terms in the log function. The problem with direct use of this expression is that Eo’ is not known precisely unless both KxCx and KyCy are known or can be reduced to zero. Either one can be easily reduced to zero by reducing the concentration of one ion to zero; however reducing the other to near zero means changing ionic strength and therefore introducing artifacts from liquid junction potentials. The problem could be solved by using an electrolyte that we are certain does not form ion pairs (e.g. K = 0); however, it is difficult to be certain that ion pairs do not form, especiallygiven the nature of the local medium around the surface ferricinium. The solution lies in comparing potentials obtained in one electrolyte to those obtained in the same electrolyte with small quantities of a second electrolyte added. In the presence of only a single electrolyte containing anion X (e.g. Cy = 0)
E,’ = Eo’ - (RT/nF)In (1 + KxCx)
(2)
Combining this with eq 1 gives
E, = E,,’ - (RT/nF)In (1 + (KyCy)/(l + KxCx))
Then, eqs 3 and 4 reduce
E, = E,,’ - (RT/nF)In ( K e g y ) (5) Equation 5 predicts the negative shift in E, as ion Y is
Kaa, M-’ 870 870 3 600 11 OOO 22 OOO
Monolayers of ferrocenylhexanethiol with n-alkanethiols. mol/cm*. b 1.0 M H&Od; V vs Ag/AgCl.
>> (1 + KxCx), or Ke&y >> 1.
to
(3)
or
Ep= E; - (RT/nF)In (1 + K e g y ) (4) where E,’ is the observed peak potential with no Y present, E, is the peak potential with Y present, and the term K y / ( l + KxCx) is an “effective”formation constant, Keel for the Fc+Y-ion pair. It is helpful to introduce the single simplifying assumption that a single ion pair is the dominant species present: e.g. that (rFc+)tot = rFc+y-,KyCy
added to solutions initially containing only ion X and provides physical meaning to the effective formation constant K e R that is derived from such data. By inserting the data from Figure 5 into eq 5, we have calculated effective formation constants for ion pairs between surface-boundferriciniumand perchlorate. Since concentrations were used instead of activities, the activity Coefficients of the anions and ferricinium are contained in K,ft. Table I contains these formation constants for monolayers containing ferrocenylhexanethiolcoadsorbed with alkanethiolsfrom 4 to 12 carbons long. The formation constant is identicalat 870M-’ for monolayersthat contain alkanethiols that are less than or equal in length to that of the ferrocenylalkane spacer; however, increasing the alkanethiol chain length by only two methylene groups above the ferrocene spacer results in an increase in the formation constant for ion pairing by a factor of four, from 870 to 3600 M-l. The formation constant continues to increase monotonically to a value of 22 000 M-l for alkanethiol coadsorbates that are 12 carbons long. Overall the formation constant increases by a factor of 25 for alkanethiol coadsorbates that are from 6 to 12 carbons long. These results confirm that as the environment surrounding the ferrocene moiety becomes more alkane-like with increasing alkanethiol coadsorbate chain length, ion pairing of ferricinium with perchlorate becomes more favored. Since perchlorate specifically ion pairs with ferricinium in these mixed monolayers, several other anions with similar chemical properties and ion-pairing tendencies were investigated in the same way as perchlorate. In the presence of 1.0 M H2S04, a monolayer composed of ferrocenylhexanethioland hexanethiol (I’= 3.0 X lo-” mol/ cm2)exhibits a negative shift in redox potential of 54 mV per decade increase in concentration of hexafluorophosphate (a representative curve for a single experiment is shown as Figure 7B). The effective ion pairing constant determined using eq 5 is 240 M-l. Trifluoroacetate and methanesulfonate were also investigated; however neither of these anions forms an ion pair as seen by the near invariance of the redox potential with changing anion concentration (Figure 7C,D). Trifluoroacetate and methanesulfonate have larger (e.g. more negative) free energies of hydration than do perchlorate and hexafluorophoaphate, which may in part explain the lesser tendency of these anions to participate in forming ion pairs. Protonation of these anions, particularly trifluoroacetate ( ~ K =A0.231, in 1.0 M HzSOl may also contribute to the absence of ion pairing. Ion pairing with PF6- is expected given that this anion has a similar ionic size and solvation free energy relative to that of C104- (C104- and PF6- are both more poorly solvated than bi~ulfate).~’These anions were also independently shown to ion pair with a mixed-valent acetylene-bridged-biferrocene in intervalence chargetransfer spectroscopic studies.31 Further evidence for ion pairing of C104-and PF6- with ferricinium is provided by ferrocene/ferricinium electron exchange experiments. Wahl and co-workerssuggested that an observed decrease in the electron exchange rate, measured by NMR line broadening in low dielectric solvents, with increasing electrolyte concentrations of PF6- and clod- may result (37) Marcus, Y. Zon Soloation; Wiley & Sons: New York, 1986; pp 107-109.
Rowe and Creager
2312 Langmuir, Vol. 7,No. 10,1991
I -3
,-
-2
-1
0
-3
-2
-1
0
L09(CAnj,,)
1 0.4 - 1
0.4
o.2 0 1
I 01 .-
01-
-t
-3
-2
0.1
-1
0
-3
-2
-1
0
Figure 7. Effect of several anions on the redox potential for electrodes coated from ethanol solutions containing (Fc)CeSH CFsC02-. and hexanethiol: (A) C104-; (B)PFa-; (C) CHsSOs-; (D) Solutions all contain 1.0 M H$04; coverage for each voltammogram is =3.0 X 10-11mol/cm2.Redox potentials were measured vs a Hg/HgpSO4/saturated KpS04 reference electrode and converted to potentials vs Ag/AgCl/saturated KC1. from ion pairing.% They analyzed their data based on this assumption and determined that the ion-pairing constant for C104-was larger than for PFe-, as we observed.
Conclusion Incorporation of ferrocenylhexanethiol into self-assembled mixed monolayers with n-alkanethiols from 4 to 12 carbons long has allowed the interfacial redox and ionpairing thermodynamics of ferrocene to be investigated. The use of n-alkanethiols as a second monolayer component enabled the control over ferrocene coverage and acted as a local medium perturbation to the redox reactivity of bound ferrocene. Increasing the length of the alkanethiols beyond that of the ferrocenylalkanethiol created an increasingly alkane-like environment around ferrocene; this caused the redox potential of ferrocene to shift to more positive potentials. The shift was explained in terms of a thermodynamic "price" that must be paid to form ferricinium in a lower dielectric alkane-like environment. An unexpected dependence of the redox potential on concentrations of c104-and PF6- indicated that ferricinium forms specific ion pairs with these ions. A mathematical model was developed to calculate effective ionpairing constants in the presence of weakly associating (38) Yang, E. 5.;Chan, Ma-S.;Wahl, A. C. J. Phys. Chem. 1980,84, 3094.
bisulfate. Ion-pairingconstants for monolayers consisting of ferrocenylhexanethiol and hexanethiol were 870 M-' for ClO4-and 240 M-l for PFs-. The ion-pairing tendencies of these ions are in qualitative agreement with previous reports involving intervalence charge transfer experimenta and solution phase electron exchange experiments as discussed in the text. Lastly, the propensity of Clod- to ion pair with ferricinium was found to increase as the local medium around ferrocene becomes more alkane-like. The effective formation constants ranged from 870 to 22 OOO M-I for monolayers composed of ferrocenylhexanethiol and alkanethiols from 6 to 12 carbons long.
Acknowledgment. This work was supported in part by a Biomedical Research Support Grant from the Division of Research Resources, National Institutes of Health (SEC). We thank Professor J. T. Hupp for bringing ref 31 to our attention. Appendix Consider the electrochemical oxidation of ferrocene coupled with two independent ion-pairing reactions of ions X and Y with ferricinium Fc F= Fc' + eFc' X- + Fc'X-
EO'
+
KX = rFc+X-/rFc+CX-
Fc' + Y- Fc'YKY= ~ F ~ + Y - / ~ F ~ + C Y The total amount of oxidized ferrocene at any time is given by + rFc+X- + rFc+YCombining this with the definitions of Kx and Ky gives rFc+
(rFc+)Tot
= r F c + ( l + KXcX + Kycy) The Nernst equation for the surface-confined ferrocene/ ferricinium redox reaction is (rFc+)Tot
+
E = E O' (RT/nF)ln ( r F c + / r F $ ) Substituting for rFc+ gives, after rearrangement E =E
O'
+ (RT/nF)In ((rFc+)Tot/rF&) -
(RT/nF)In (1+ KxC,
+ KyCy)
We can assume that ferrocene itself does not ion pair, e.g., that (rFcO)Tot = rFcO. We then note that, at E, (for a symmetric wave), (rFc+)Tot = rFcO. Then the second term on the right side of the equation above goes to zero, and
E, = E O' - (RT/nF)In (1+ KxCx + KyC,) Registry No. Au, 7440-57-5;ferrocene, 102-54-5; ferrocenylhexanethiol, 134029-92-8; ferrocenyl 5-bromopentyl ketone, 57640-76-3; ferrocenyl-6-bromohexane,136237-36-0; n-butanethiol, 109-79-5;n-hexanethiol, 111-31-9;n-octanethiol, 111-88-6; n-decanethiol, 143-10-2.
