Langmuir 1993,9, 3668-3675
3668
Self -Assembled Monolayers in Organic Solvents: Electrochemistry at Alkanethiolate-Coated Gold in Propylene Carbonate Kimberly A. Groat and Stephen E. Creager’ Department of Chemistry, Indiana University, Bloomington, Indiana 47405 Received September 20,199P
Alkanethiolate monolayers with excellent barrier properties can be formed and maintained on gold electrodes in the nonaqueous solvent propylene carbonate. The layers are most stable when alkanethiol is present in the propylene carbonate solution and when the potential is held between +0.50 and -0.70 V vs Ag/AgCl/saturated KC1. The ability to work in a nonaqueous solvent has enabled study of several redox-active probe molecules that have not previously been studied at alkanethiolate-coatedelectrodes. Specifically,studies on three ferrocenederivatives (ferrocene,decamethylferrocene,and [(trimethylaminolmethyl]ferrocene) and three organic redox mediators (tetracyanoquinodimethane,N,N,N”-tetramethylphenylenediamine, and NP-dimethylviologen) have been pursued. Two complex metal ions (hexacyanoferrate(II1) and hexaa”ineruthenium(II1)) were also studied to facilitate comparison with earlier work in water. A weak correlation exists between the rate of interfacial electron transfer at the coated electrodes and the homogeneous electron self-exchange rates of the redox-active probe molecules. The correlation is strongest within the series of structurally homologous ferrocene derivatives and is weakest when comparing probe molecules of very different structure. There is no apparent correlation between electrochemical reactivity and molecular size, suggesting that size exclusion is not a determining factor for electrochemistryat alkanethiolate-coatedelectrodes. There is however a correlationwithmolecular charge, suggesting that electrostatic effects may be important.
Introduction Molecular self-assembly onto metal electrode surfaces has emerged as an attractive means of preparing chemically designed surfaces for use in fundamental research on interfacial charge transfer and for practical applications in electroanalysis. Alkanethiol-derived monolayers selfassembled onto gold are particularly well behaved and have been the subject of much recent research.lv2 Of particular importance is the capacity of such monolayers to serve as barrier layers toward approach of solutionphase species to an electrode surface. Blocking the approach of electroactive solutes forces electron transfer to occur over a relatively long distance across the monolayer; this is of fundamental interest with respect to the mechanism of long-range electron transfer.sg Blocking the approach of solvent and electrolyte ions brings about a dramatic decrease in the interfacial capacitance and therefore the capacitive charging current;sJOthis in turn provides advantages in electr~analysis.~~-~~ Sophisticated strategies for forcing electron transfer to occur at molecularly designed “gate sites” in mixed monolayers are being *Abstract published in Advance ACS Abstracts, November 15, 1993. (1) Ulman,A. An Introduction to Ultrathin Organic Films. From
Lagmuir-Blodgett to Self-Assembly.;Academic Press, Inc.: San Diego, CA, 1991; p 442. (2) Dubois, L. H.; Nuzzo, R. G. Annu. Rev. Phys. Chem. 1992,43,437. (3) Becka, A. M.; Miller, C. J. J. Phys. Chem. 1992,96, 2657. (4) Miller, C.; Gratzel, M. J. Phys. Chem. 1991, 95, 5225. (5) Miller, C.; Cuendet, P.; Gratzel, M. J. Phys. Chem. 1991,95,877. Avery, S.; Lynch, M.; Furtach, T. Langmuir 1987, (6) Finklea, H.0.; 3, 409. (7) Morieaki, H.; Niehikawa, A,; Ono, H.; Yazawa, K. J. Electrochem. SOC. 1990,137, 2759. (8) Porter, M. D.; Bright, T. B.; Allara, D. L.; Chidsey, C. E. D. J.Am. Chem. SOC. 1987,109,3559. (9) Demoz, A.; Harrison, D. J. Langmuir 1993,9, 1046. (10) Chidsey, C. E.D.; Loiacono, D. N. Langmuir 1990,6,682. (11) Wang, J.; Wu, H.; Angnes, L. Anal. Chem. 1993,155, 1893. (12) Malem, F.; Mandler, D. Anal. Chem. 1993, 66,37. (13) Quan, C.; Brajter-Toth, A. Anal. Chem. 1992,64,1998. (14) Sabatani, E.; Rubinstein, 1. J. Phys. Chem. 1987, 91, 6663.
developed as a means of introducing chemical selectivity into interfacial electron-transfer reactions.lbl8 An apparent limitation of the gold/alkanethiol chemistry is that the alkanethiolate monolayers appear to be wellbehaved only in water. Scattered reports regarding the electrochemistry of immobilized electroactive adsorbates in nonaqueous solvents have appeared;lSz2 however attempts to use alkanethiolate monolayers as barrier layers in nonaqueous solvents have for the most part been unsuccessful,presumably because of monolayer instability. This is unfortunate, since many of the most interesting and potentially revealingredox-activeprobe molecules that one might like to investigate at a blocked electrode are incompatible with aqueous media, because of either solubility or chemical stability considerations. Also, varying the solvent is an important strategy in research on electron transfer, since contemporary theory makes specific predictions regarding how such reactions should be affected by changes in solvent properties.23-w We report herein the results of our study on the barrier properties of alkanethiolate monolayers on gold in propylene carbonate. Monolayers with excellent barrier properties can be formed and maintained in this solvent as long as a small quantity of alkanethiol is present in the contacting solution and the applied potential is kept within (15) Bilewicz, R.; Majda, M. Langmuir 1991, 7, 2794. 1991,113,5464. (16) Bilewicz, R.; Majda, M. J. Am. Chem. SOC. (17) Chailapakul, 0.;Crooks, R. M. Langmuir 1993,9,884. (18) Steinberg, S.; Rubinatein, I. Langmuir 1992,8, 1183. (19) Hickman,J.J.;Laibinie,P.E.;Auerbach,D.I.;Zou,C.F.;Gardner, T. J.; Whitesides, G. M.; Wrighton, M. S . Langmuir 1992,8, 357. (20) Hickman, J. J.; Ofer, D.; Zou, C.; Wrighton, M. 5.;Laibinis, P. E.; Whitesides, G. M. J.Am. Chem. SOC. 1991,113,1128. (21) Curtin, L. S.; Peck,5.R.; Tender, L. M.; Murray, R. W.; Rowe, G. K.; Creager, S. E.A d . Chem. 1993,66, 386. (22) Tsutaumi, H.;Furumoto, S.; Morita, M.; Matauda, Y. J. Electrochem. SOC. 1992,139,1522. (23) Rips, I.; Jortner, J. Chem. Phys. Lett. 1987, 133, 411. (24) Rips, I.; Jortner, J. J. Chem. Phys. 1987,87, 2090. (25) McManis, G. E.; Nielson, R. M.; Gochev, A.; Weaver, M. J. J.Am. Chem. SOC. 1989,111,5533. (26) Fawcett, W. R.; Blum, L. Chem. Phys. Lett. 1991, 187, 173.
0743-7463/93/2409-3668$04.00/00 1993 American Chemical Society
Self-Assembled Monolayers in Organic Solvents chart 1
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DMFc
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a specied range. For a propylene carbonate solution containing 0.1 M tetraethylammonium tetrafluoroborate and as little as 5 p M dodecanethiol, that range is +0.50 to -0.70 V vs Ag/AgCl/saturated KC1. We have investigated the electrochemistry of eight common redox-active probe molecules at alkanethiolate-coated gold in propylene carbonate, including three substituted ferrocene derivatives (ferrocene, decamethylferrocene, and [(trimethy1amino)methyllferrocene),three organic redox mediators (tetracyanoquinodimethane,NJV,ZV'JV'-tetramethylphenylenediamine, and N,N'-dimethylviologen), and two complex metal ions (hexacyanoferrate(II1)and hexaammineruthenium(1II)). Structures for several of these redox-active probe molecules are presented in Chart 1. In general terms, we find that redox probe molecules that exhibit fast homogeneous electron self-exchange kinetics also exhibit relatively rapid interfacial electrontransfer kinetics at a coated electrode. Within the ferrocene series this correlation is obeyed very well; when comparisons are made among structurally unrelated probe molecules, however, the correlation is much poorer. A common element in our observations is that electrically neutral molecules consistently exhibit more facile interfacial redox kinetics at coated electrodes than do charged molecules, especially highly charged molecules. Possible implications of these observations regarding adsorption, partitioning, and reaction at microscopic and macroscopic defect sites in the monolayers are discussed. Experimental Section Electrodes and Cells. Gold electrodes were constructed by sealing annealed (lo00 "C) gold wires in an epoxy resin (EPON 825, Shell) that was cured using n-phenylenediamine (Aldrich) as hardening agent. The resin and hardening agent were combined in a 100.13weight ratio at approximately 65 "C, drawn into a length of 0.25 in. diameter i.d. polyethylene tubing containing a stub of gold wire soldered onto a braided copper wire, and cured at approximately 70 "C overnight. A gold disk was then exposed by sanding the end of the assembly, and the resulting electrode was smoothed by polishing with an alumina slurry. The macroscopicarea was estimatedvisually to be 0.0105 cm2. This electrode geometryhas proved much more convenient thanthat described in our earlier publication on polycrystalline bulk gold electrodesn (eg., unehrouded planar squares) since these electrodes are easier to handle and can be immersed in solution, thereby eliminating the need for noncontact and/or O-ring cells. The electrochemicalcellconsisted of a 25-mL three-neckpearshaped flask with 14/20 standard taper ground-glass joints to hold auxiliary,reference,and working electrodes. Electrochem(27) Creager, 5. E.; Hockett, L.A.; &we, G. K.Langmuir 1992,8,854.
Langmuir, Vol. 9, No. 12,1993 3669
ical data were acquired and are reported versus a Ag/AgCV saturated KCl reference electrode which was isolated from the main cell compartment by a salt bridge. All experimentsutilized a coiled platinum wire auxiliary electrode,and all solutions were spargedwith solvent-saturatednitrogen prior to data acquisition. Immediately before use the electrodes were polished on felt using aqueous slurries of progressively finer alumina (25, 1,0.3 pm), sonicated in soapy water to remove the alumina,and etched using dilute aqua regia (31:4 concentrated HCkconcentrated HN0a:HzO) in a quartz test tube for approximately 1 min. Immediately after being etched, the electrodes were rinsed with water and ethanoland then immersed in a 20 mM solution of the appropriate alkanethiol in ethanol for no less than 8 h to form the desired coating. This procedure has been shown to yield self-assembled monolayer films of very high integrity." Instrumentation and Methods. Voltammetric data were acquired using a single channel of a Pine Model RDE-4 bipotentioatat and recorded on an X-Yrecorder (YokagawaModel 3023or 3025). Capacitance data were acquired using an EG&G PAR Model 362 potentioatat with a 20 Hz/20 mV peak-to-peak trianglewave superimposedon the dc bias potential;capacitance was calculated from the resulting square-wavecharging current signal displayed on a Tektronix Model TDS-320 digital oscilloscope. GC-MS data were acquired using a Hewlett-Packard Model 5890series 11gas chromatograph(SupelcoSPB-5column, 60 m X 0.25 mm) coupled with a Hewlett-Packard Model 5971 mass-selective detector. Propylene carbonate solutions to be analyzed by GC-MS were extracted with pentane, and the pentane volume was reduced to approximately 10 fiL prior to analysis. The platinum auxiliary electrode was kept behind a separate frit in these experimentsto prevent thiol readsorption. Materials. Propylene carbonate (Baker) was purified by vacuum distillationfrompotaaaium permangauate.28 Tetraethylammonium tetrafluoroborate (Aldrich) was recrystallized twice from methanovwater. Absolute ethanol (Midwest Grain Producta), dodecanethiol (Aldrich),ferrocene (Fc, Strem),and decamethylferrocene (DMFc, Strem) were wed as received. Octanethioland hexanethiol(Aldrich)were passed through activated alumina before use. [(Trimethylamino)methyllferroceneh e d u orophosphate (TMAMFc+PFe-)was prepared by condensation of [(dimethylamino)methyllferrocene (Aldrich) with methyl iodide, followed by precipitation with potassium hexafluorophosphate in water. Femceniumhexafluorophoaphate (Fc+PFe-) was prepared by oxidation of ferrocenewith concentratedsulfuric acid, followed by precipitation with potassium heduorophosphate in water.a Tetracyanoquinodimethane (TCNQ,Aldrich) and NJvJv'Jv'-tetramethylphenylenediamine (TMPD,Aldrich) were purified by vacuum sublimation prior to use. N&'Dimethylviologen bie(heduorophosphate) (W+(PFe-)a)was prepared from the corresponding dichloride (Aldrich) by precipitation using potassium hexafluorophoaphate in water. Hexaammineruthenium(II1) tris(hexafluoropho8phate) (Ru(NHa)$+(PFe-)s) was prepared from the trichloride salt
(Aldrich)by precipitation with potassium hexafluorophosphate in water. Tetrabutylammonium hexacyanoferrate(II1) (((C&)fi+)fle(CN)ea) was prepared by exposureof an aqueous solutionof potassiumhexacyanoferrate(III1to a dichloromethane solutionof tetrabutylammonium chloride in a separatory funneLm The tetrabutylammonium hexacyanoferrate(II1) product was precipitated from dichloromethane by addition of diethyl ether. Results and Discussion I. Alkanethiolate Monolayers on Gold in Propylene Carbonate. A primary goal of this study was to extend the use of alkanethiolate blocking layers on gold to nonaqueous solvents. Propylene carbonate is an excellent choice as an alternative solvent; it is quite polar and therefore easily dissolves electrolyte, electrolyte solutions in it are highly conductive, and it is rigorously aprotic (28) Perrin, D. D.;Armarego, W.L. F. Purification of Laboratory Chemicals, 3rd ed.; Pergamon Press: Oxford, U.K., 1988. (29) Yang, E. 5.;Chan, MA.;Wahl, A. C. J. Phys. Chem. 1980,84, 3094. (30)Des., B.; Cerlin, R.; Osteryoung, R. A. Znorg. Chem. 1989,28,421.
Groat and Creager
3670 Langmuir, Vol. 9, No. 12, 1993
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Figure 1. Cyclicvoltammetryof ferrocene(0.2 mM) inpropylene carbonate containing 0.1 M tetraethylammonium tetrafluoroborate. All scans 100 mV sec-1. (A) Electrode immersed in dodecanethiol solution in ethanol prior to use. The arrows indicate the shift in peak position with continued scanning. (B) Bare gold electrode, 50 pM dodecanethiol added to solution immediately prior to start of scanning. Arrows again indicate shifts in peak potential with continued scanning. (C)Fresh voltammogram taken immediately after (B) using the same electrode. (D)The electrode was reexposed to dodecanethiol solution in ethanol, and reexamined in electrolyte solution from (B)containing 50 p M dodecanethiol. with a wide potential window that makes it suitable for experimentswith a broad range of redox-active molecules. Our initial attempts at using alkanethiolate monolayers as blocking layers in propylene carbonate were unsuccessful, due to an apparent tendency of alkanethiolates (or derivatives thereof) to come off the surface. Figure 1A illustrates this problem at dodecanethiolate-coated gold in propylene carbonate containing 0.2 mM ferrocene and 0.1 M tetraethylammonium tetrafluoroborate. An irreversible voltammetric wave with a positive shift in the peak potential (Epk = +0.62 v at 100 mV s-l, compared with the reversible value of +0.47 V) is initially observed at this electrode, indicating that the alkenthiolate monolayer impedes electron transfer between the electrode and ferrocene. As scanning continues, however, the oxidative peak shifts negative and a rereduction peak grows in until the voltammogram eventually approaches the familiar form characteristic of reversible (Le,, rapid) electron transfer at room temperature (e.g., 60 mV peak splitting). We postulate that the alkenthiolate monolayer was gradually coming off the electrode and that electron transfer was occurring at defect sites that were left behindee To alleviate this situation, we decided to add a small amount of dodecanethiol (50 p M ) to the propylene carbonate electrolyte solution. Given the strong affinity of thiols for gold,31we anticipated that even this small concentration of alkenthiol might serve to replace any material that leaves the electrode surface, thereby keeping the blocking capacity of the monolayer intact. Evidence that this procedure works is seen in Figure 1B-D. If 50 pM of dodecanethiol is addedto the solution and scanning immediately begun at an initially uncoated gold electrode, then the voltammetry for ferrocene oxidationireduction slowly changes from the reversible form back to the (31)Nuzzo, R.G.;%garski, B. R.; Dubia, L. H. J . Am. Chem. SOC.
1987,109, 733.
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Figure2. Background voltammetricscam at dodecanethiolatecoated gold in propylene carbonate: top, no thiol in solution; bottom, solution contains 50 p M dodecanethiol. Scan rate 100 mV 8-l.
irreversible form observed initially at a coated electrode. This behavior supports our postulate that loss of alkanethiolates creates defect sites where electron transfer is rapid but that those defects can be "plugged up" by adding a small quantity of alkanethiol to the solution. When the same electrode is polished, etched, reimmersed in dodecanethiol solution in ethanol, and placed in an electrolyte solution containing 0.2 mM ferrocene and 50 p M dodecanethiol, the characteristic irreversible voltammetry is again observed but is now entirely stable for at least the duration of an experiment (typically a few hours). In the study of such layers on electrodes, the monolayer stability as a function of potential is of particular importance. It is known for example that in water, excursions to extreme negative or positive potentials result in irreversible loss of alkanethiolate from a coated electrode by electrochemical oxidatiodreduction processes involving the thiolate Figure 2 illustrates that the barrier properties of alkanethiolate monolayers on gold in propylene carbonate are also compromised beyond certain potential limits. At the top of Figure 2 is a set of cyclic voltammograms for a dodecanethiolate-coated electrode in propylene carbonate that does not contain any alkanethiol. The background current is initially quite low at thia electrodebut begins to rise with continued scanning to positive potentials. The bottom voltammogramswere acquired in a similar manner in a solution containing 50 pM dodecanethiol; the background current is clearly much lower and s t a y relatively constant with scanning, indicating that the monolayer barrier properties are improved by having alkanethiol present in solution. The voltammograms still exhibit a rising background at potentials beyond +0.5 V, however, indicating that the barrier properties of the monolayer are at least partially compromised at these potentials. A more quantitative indicator of how the background current varies with potential is the interfacial capacitance. An ideal barrier layer should exhibit a very low capacitance that varies minimally with potential. Figure 3 presents data on capacitance us potential, measured by the trianglewave method, at a dodecanethiolate-coated electrode in propylene carbonate containing dodecanethiol. The ca(32) W e l d , M. M.; Popenoe, D. D.; Deinhammer, R. 5.; Lamp, B. D.; Chung, C.; Porter, M . D. Langmuir 1991, 7, 2887.
Self-Assembled Monolayers in Organic Solvents
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pacitance exhibits a plateau value of approximately 3 pF cm-2 between +0.50 and 4.70 V. This plateau value is quite low relative to that expected at a typical bare metal electrode (>20 pF cm-% Outside of this potential range the capacitance begins to rise, indicating that the barrier properties of the monolayer with respect to solvent and/ or electrolyte penetration have been breached. We postulate that a potential-induced desorption of alkanethiolate is responsible for this behavior. In support of this postulate, we note that if an alkanethiolate-coated electrode is immersed in propylene carbonate solution at a potential within the capacitance plateau (e.g., 0.0 V) and then held at that potential for several minutes, the monolayer barrier properties are nearly identical to those of a freshly coated electrode immediatelyafter immersion. Monolayer barrier properties are compromised only after excursion to potentials outside of the capacitance plateau. For comparison, we have included in Figure 3 capacitance data acquired at an identically prepared electrode in aqueous perchloric acid solution. The general behavior in the two solvents in similar, although the plateau capacitance is approximately one-third lower and the positive potential limit slightly less positive in water than in propylene carbonate. We hypothesize that in propylene carbonate, solvent molecules can partially penetrate the monolayer to a greater extent than in water, such that the monolayer becomes more polarizable and/or has a smaller effective thickness as a barrier layer toward solvent, thereby increasing the capacitance. Experiments using different concentrations of alkanethio1 in the electrolyte revealed similar behavior in both the ferroceneblocking and interfacial capacitance experiments for alkanethiol concentrations as low as 5 pM and as high as 5 mM. It is significant that the use of fairly high concentrations of alkanethiol did not yield a continued shift to more positive peak potentials for ferrocene oxidation or a continued decrease in the plateau capacitance. This suggests that multilayer structures that could block the electrode surface more thoroughly than a monolayer are not formed. In support of our postulate of potential-induced alkanethiolate desorption, we have performed experiments in which an electrode of relatively large macroscopic area
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coated in a dodecanethiol solution in ethanol, immersed (after rinsing) at 0.0 V into a small volume (around 1 mL) of propylene carbonate containing electrolyte, and then emersed at a desired potential on or off the capacitance plateau. After repeating the coating/immersion/emersion steps 9 times, we then analyzed the propylene carbonate solution (after workup with pentane as described in the Experimental Section) for thiolate-derived products by gas chromatography-mass spectrometry. Our hypothesis was that emersion at a fairly positive potential (we used +0.8 V) should result in accumulation of thiolate-derived species in solution, whereas emersion at a potential in the plateau region (we used 0.0 V) would not. Figure 4 illustrates that this is indeed the case; a total ion chromatogram corresponding to emersion at +0.8 V (top) yields a peak at a retention time identical to that obtained for an authentic sample of dodecanethiol,whereas the parallel chromatogram corresponding to emersion at 0.0 V shows no such peak. The full mass spectrum (bottom) Corresponding to this peak is also nearly identical to that obtained from an authentic sample of dodecanethiol. Control experiments with an authentic sample of dodecyl disulfide worked up from propylene carbonate in a fashion similar to that used in the electrochemical desorption experiments yielded a peak at much longer retention times with a mass spectrum similar to that of authentic dodecyl disulfide, indicating that potentialinduced desorption of alkanethiolate from the electrode does not yield disulfide as the primary product. Apparently, desorption of alkanethiolate occurs in concert with proton and/or hydrogen-atom transfer to yield the thiol as the primary product. One possible mechanism that would account for this is that the alkanethiolate is desorbed as a radical species, and that subsequent reactions with solvent, electrolyte, and/or trace water produce the thiol. No other thiolate-derived products besides the thiol were observed in our experimenta; however we cannot rule out the presence of hydrophilic and/or nonvolatile products
Groat and Creager
3672 Langmuir, Vol. 9, No. 12, 1993
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Figure8. Voltammetry of propylenecarbonate solutionscontaining0.1 Mtetraethylammonium tetrafluoroborate and variousferrocene derivatives. In each case, the top voltammogram was acquired at a bare gold electrode with no thiol in solution, and the bottom one at a dodecanethiolate-coated electrode with 50 pM dodecanethiol present in solution. All scans 100 mV/s. Left,solution contains 0.2 mM Fc; center, solution contains 0.2 mM DMFc; right, solution contains 0.2 mM TMAMF’c+PFs-.
that either do not partition into pentane or are too nonvolatile to be analyzed by gas chromatography. More definitive conclusions regarding the mechanism of the potential-induced alkanethiolate desorption must await the results of further experimental work in which all thiolate-derived products are accounted for. 11. Redox Reactions of Some Benchmark W o x Probe Molecules. A. Ferroceneand Ferrocene Derivatives. The ferrocene/ferricenium redox system is ideal for fundamental studies of interfacial electron transfer, in part because of the ease with which structural diversity can be achieved by making simple functional group substitutions onto the cyclopentadienyl rings. Another of our goals in this study was to exploit this structural diversity to learn how the blocking of interfacial redox reactions at alkanethiolate-coated electrodes depends on the structure of the redox-active probe molecule. We have pursued this idea using a series of three ferrocene derivatives: ferrocene (Fc); decamethylferrocene (DMFc); and [(trimethy1amino)methyllferrocenehexafluorophosphate (TMAMFc). DMFc is unique in that it is simply an expanded ferrocene; the molecular volumes of DMFc and Fc differ by a factor of 2.21 (451 vs 204 A3 from crystallographic datas39, which suggests that their effective radii will differ by a factor of (2.21)’/3 = 1.30. TMAMFc is unique in that it is cationic in both of its redox states. By studying the two structural derivatives DMFc and TMAMFc, we therefore expect to gain b i g h t into how two important factors, size and charge, affect the capacity of the monolayer to block electron transfer. Figure 5 illustrates the capacity of a dodecanethiolate monolayer on gold to inhibit the redox reactions of each of these three ferrocene derivatives. Significant positive shiftsin oxidative peak potential relative to the reversible value are observed for all three compounds. (Allscans were at 100 mV However, the shifts for ferrocene (170 mV) and for [(trimethylamino)methyllferrocene(270 mV) are clearly much larger than that for decamethylferrocene (15mV). The differences in voltammetric response among these three molecules were observed for many different electrodes in many individual experimentsand are entirely reproducible. In one instance we examined a single (33) Freyberg, D.P.;Robbins, J. R;Raymond, K.N.;Smart, J. C.J. Am. Chem. Soc. 1979,101,092. (34)Seiler, P.;Dunitz, J. D. Acto Crystollogr. 1979, B95, 1068.
electrode in three separate solutions, each containing one of the three indicated ferrocene derivatives (in addition to electrolyte and alkanethiol); the differences in voltammetric response were observed even when the same electrode was used for all three measurements,c o n f i i g that they are not due to random fluctuations in the quality of the monolayers. The observation that DMFc+/Ois the most facile of the three ferrocene reactions is contrary to our initial expectation that a larger molecule would be blocked better than a smaller one since it should be less able to penetrate into microscopic structural defects in the monolayer. If penetration at defect sites was the dominant factor controlling the apparent electron-transfer rates, then it follows that the larger molecule should exhibit correspondingly slower redox kinetics. This is clearly not the case. We conclude that, at the very least, relative degrees of penetration of these ferrocene molecules into structural defect sites are playing only a minor role in determining the apparent electron-transfer rates at the alkanethiolcoated electrodes. B. Tetracyanoquinodimethane,Tetramethylphenylenediamine,and Dimethylviologen. Another strategy for studying the effect of probe molecule structure, particularly charge, on the capacity of a monolayer to block electron transfer is to utilize molecules that can undergo more thanone redox reaction. This strategy is particularly powerful when the various redox statesof a probe molecule all have very similar internal structures; this condition corresponds to outer-sphere electron transfer, for which simple continuum theories predict that the intrinsic activation parameters should be similar for each redox reaction. Three common probe molecules that meet these requirements are tetracyanoquinodimethane (TCNQ), NJV’-dimethylviologen (MV2+),and NJVJV’JV-tetramethylphenylenediamine (TMPD). TCNQ and M V + undergo two sequential one-electron reductions, and TMPD undergoes two sequential one-electron oxidations. In each case the benzenoid skeleton is preserved such that the internal bonding is very similar in all the states involved in the relevant redox reactions. Figures 6 8 illustrate how the electrochemistry of these molecules a t gold is affected by the presence of a dodecanethiolate monolayer. The behavior of TCNQ is particularly illustrative;the f i t redox reaction, TCNQOi-,
Langmuir, Vol. 9, No.12,1993 3613
Self-Assembled Monolayers in Organic Solvents
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-0.2
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Figure 8. Voltammetry of MV2+(PFe-)2,approximately mM concentration, in propylene carbonate: top, bare gold, S = 0.1 pA, no thiol in solution;bottom, dodecanethiol-coatedgold, S = 0.06 PA, solution contains 50 mM dodecanethiol. Other conditions identical to Figure 5. appears completely reversible at the coated electrode, whereas the peak potential for the second redox reaction, TCNQ-/%,has been shifted negative by almost 200 mV at the coated electrode relative to the reversible value. The first redox reaction of TMPD, TMPD+/O,is not reversible but is still readily observed (E,, is shifted positive by approximately 120 mV relative to the reversible value), whereas the peak potential corresponding to the second redox reaction, TMPD2+/+,is apparently shifted so far positive that it is not observed within the potential range examined. Both these results are indicative of a tendency for redox reactions of charged molecules to be blocked
more effectively than those of neutral molecules. The electrochemistry of MV2+ at a coated electrode strongly supports this hypothesis; the peak potential correaponding to reduction of MV2+is shifted negative by over 500 mV relative to the reversible value, by far the largest shift among these three molecules. The observed wave for MV2+/+is in fact nearly atop that expected for MV+/o, which precludes examination of that redox reaction. Altogether, these data strongly suggest that interfacial electron transfer at a monolayer-coated electrode is inhibited more effectively for charged molecules than for neutral molecules. C. Hexaammineruthenium and Hexacyanoferrate. It is informative to compare our results in propylene carbonate with the results of prior work at alkanethiolatecoated electrodes in water. Two redox systems that have been extensively studied in water, both at uncoated%* and ~ ~ a t e d ~metal * ~electrodes, * ~ J ~are~the ~ hexa~ ~ ~ ammineruthenium and hexacyanoferrate systems. The high chemical stability of these two complex ions, particularly toward ligand substitution, and the wealth of information available on their redox chemistrymake them excellent benchmark systems for fundamental studies on interfacial redox chemistry. Relatively little work has been reported for these two systems in nonaqueous solvents, no doubt due to solubility considerations. However, by judicious choice of counterion, it is possible to get both hexaa"ineruthenium(II1) and hexacyanoferrate(II1) into propylene carbonate solution. Figures 9 and 10illustrate how the electrochemistry of these molecules in propylene carbonate is affected by the presence of an alkanethiolate monolayer on the electrode surface. Our initial studies at dodecanethiolatecoated gold showed essentially complete blocking for these molecules, i.e. no voltammetric waves could be observed within the potential range examined. Hence, in an attempt to bring the kinetics into a range where they could be more easily examined, we used two shorter alkanethiols, hexanethiol and octanethiol,to coat the electrodes. Figure 9 illustrates that well-behaved voltammetry could be (35) Peter, L. M.; Dun,W.; Bindra, P.; Gerisher, H. J. Electrochem. Chem. 1976, 71, 31. (36) &Met& T.; Weaver, M. J. A d . Chem. 1984,66, 1444. (37) Sabatani, E.; Rubinstein, I.; Maoz, R.;Sagiv, J. J. Electrochem. Chem. 1987,219, 365.
Groat and Creager
3674 Langmuir, Vol. 9, No. 12, 1993
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Figure 10. Voltammetry of Ru(NH&(PF,&, concentration approximately mM, in propylene carbonate: top, bare gold electrode, S = 0.1 pA, no thiol in solution; bottom, octanethiolatecoated gold, S = 0.06pA, solution contains 50 pM octanethiol. Other conditions identical t o Figure 5. Table 1. Selected Electron-Transfer Kinetics Data for Some Benchmark Redox Reactions molecule and couple TCNQOITMPD+/O W+/O
TMPD2+/+ MV2+/+ DMFcO/+ FcO/+ TMAMFc2+/+ Ru(NH&'+/'+ RU(NH~)P/~+ (corrected) Fe(CN)6s/C
khmo(M-I 3-4 x 109 1.0-1.2 x log 1.0 x 108 6.7 x 107
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'Acetonitrile, refs 42 and 43. bAcetonitrile, refs 44 and 45. Dimethylformamide, ref 46. Acetonitrile, refs 25, 29, and 47. e Acetone, ref 48. f Acetonitrile, refs 25,29,47,49, and 50.8 Acetonitrile, refs 49 and 51-54. h Acetonitrile, refs 29 and 55. Water, refs 56-58. j Water, refs 36,53, and 59. Water, refs 36 and 59. Water, refs 60 and 61. Water, ref 35.
obtained for the Fe(CN)&/& redox reaction at a hexanethiolate-coated electrode. The peak potential for Fe(CN)& reduction is shifted 380 mV relative to the reversible value, indicating that even this relatively thin alkanethiolate layer is effective in blocking electron transfer. (Note that the redox potential for Fe(CN)&/& is dramatically shifted in propylene carbonate relative to water. This effect, which is presumably due to both the different solvating abilities of propylene carbonate and water toward polyvalent anions and to the hydrophobic nature of the tetraalkylammonium cations, has been previously observed.30) Similar behavior is seen for the R u ( N H ~ ) ~ ~redox + / ~ +reaction a t octanethiolate-coated gold; the peak potential for Ru(NH3)e2+oxidationis shifted 690 mV relative to the reversible value, again indicating very effective blocking of the interfacial redox reaction. The generaltrend is that these redox reactions are inhibited much more effectively by the alkenthiolate monolayers than are any of the reactions discussed above, which provides further support for our postulate that interfacial redox reactions of highly charged probe molecules are blocked most effectively. D. Mechanisms of Interfacial Electron Transfer at Blocked Electrodes. In considering some possible causes for the observations presented above, it is helpful to present a compilation of some kinetics data on electron self-exchangereactions involvingthe relevant redox probe molecules. Table 1presents such a compilation, including rates for both homogeneous and heterogeneous electron transfer where available. The solvent is in most cases either acetonitrile or water (see footnotes). Homogeneous electron self-exchangerate constants are typically obtained
from ESR or NMR line-broadening or isotopic exchange data; heterogeneous electron-transfer rate constants are for bare metal electrodes, typically mercury, gold, or platinum. Rate constants are uncorrected for electrostatic double-layer or work-term effects except as indicated. A common trend in our data is that redox couples that exhibit fast homogeneous electron self-exchange kinetics (e.g. TCNQOI-) also exhibit fast interfacialelectron-transfer kinetics at alkanethiolate-coated electrodes. Similarly, couples that exhibit relatively slow homogeneous kinetics (e.g. Fe(CN)s3-/& and R u ( N H ~ ) ~ ~ + also /~+ exhibit ) slow interfacial kinetics at coated electrodes. Within the ferrocene series this trend is followed with excellent integrity, despite the fact that the homogeneous selfexchange rate constants for the three ferrocene derivatives differ by only a factor of 1&20. It is tempting to conclude from this correlation that interfacial electmn-transfer rates at alkanethiolate-coated electrodes are determined largely by the intrinsic activation parameters associated with homogeneous electron self-exchange. Upon closer examination, however, the correlation is poorer, especiallywhen structurally unrelated redox couples are compared with each other. For example, reduction of MV2+is inhibited much more effectively at a coated electrode than is oxidation of any of the three ferrocene derivatives, yet the homogeneous electron self-exchange rate constant for MV2+/+is higher than that for any of the ferrocene redox reactions. Also, ferrocene oxidation at a coated electrode appears to be inhibited to nearly the same extent as is TMPD oxidation, yet the homogeneous electron selfexchange rates for these two redox reactions differ by a factor of over 100. It seems likely that something more than simple longrange electron transfer, with the alkanethiolate monolayer acting as an inert barrier, is occurring. One possibility is that adsorption of probe molecules onto the monolayer surface occurs to different degrees for each system. A favorable, adsorptive interaction with the local environment at the electrode surface can decrease the effective activation energy for electron transfer just as an unfavorable, repulsive interaction can increase it.% This is consistent with neutral molecules being more rapidly oxidized/reduced than charged molecules, since neutral molecules are more likely to be adsorbed onto a hydrophobic monolayer. Another possibility is that neutral molecules can more closely approach the electrode surface than can charged molecules, thereby allowing them to achieve stronger electronic overlap with the electrode. Close approach could be achieved by the more intimate contact associated with adsorption of a probe molecule onto the monolayer or by partial penetration of a probe molecule into the monolayer. This effect is distinct from the adsorption effect discussed above in that adsorption by itself affects primarily the activation energy for the reaction, whereas the distance of closest approach affects primarily the adiabaticity, Le. the preexponential factor. In principle, adsorption and distance effects could be distinguished by independent measurement of adsorption energies from solution and/or from temperature-dependent kinetics data which yield activation parameters for the reaction. The possibility of reaction at defect sites in the monolayer must also be considered. It is a simple fact that microscopic defects are undoubtedly present in even the most well-formed of these monolayers. Topographic (38)Weaver, M. J. InElectron transjer and electrochemicalreactions. Photochemical and other energized processes; Zuckerman, J. J., Ed.; VCH Publishers, Inc.: Deerfield Beach, FL, 1986; Vol. 15; p 147.
Self-Assembled Monolayers in Organic Solvents
substrate defects such as step edges and crystal grain boundaries are inevitable, as are domain boundaries between regions of differing chain tilt or with different registry of thiolate groups with the underlying gold. STM imaging of our etched gold electrodes has confirmed that macroscopic structure is present,2I and recent X-ray and helium atom scattering data confirm the presence of domain defects in alkanethiolate monolayers on a Au(111)single ~ r y s t a l . ~Raman ~ ? ~ spectroscopic data furthermore indicate that a small but significant quantity of the gauche conformation of the alkyl chains exists in alkanethiolate monolayers on gold;41these conformers can exist only at microscopic defects sites since they would disrupt the packing in a well-ordered, all-trans monolayer. Defects can also be formed at certain potentials via electrochemical processes, as discussed above. Attempts to totally eliminate these microscopic defects are probably an exercise in futility; the most one can hope for is to limit their number and size such that they have only a minor effect and that the average behavior is that of a blocked electrode. Regarding the role of defects in our experiments, it seems unlikely that macroscopic defects, i.e pinholes with large regions of bare metal exposed, are playing a dominant role, since if they were, the apparent interfacial electrontransfer rates at coated electrodes should correlate better with the tabulated homogeneous and heterogeneous rates. If electron transfer were occurring at microscopic defect sites, i.e. defects of molecular dimensions, then it would be expected that neutral molecules would react faster since they should be better able to penetrate the defects. The observation that the larger DMFc is oxidized faster than the smaller Fc argues strongly against a size-selection effect (39) Fenter, P.;Eisenberger, P.; Liang, K. S.Phys.Reu. Lett. 1993,70, 2447. (40) Camillone, N.; Chidsey, C. E. D.; Eisenberger, P.; Fenter, P.; Li, J.; Liang, K. S.; Liu, G.-Y.; Scoles, G. J. Chem. Phys. 1993, 99, 744. (41) Bryant, M. A.; Pemberton, J. E. J. Am. Chem. SOC.1991, 113, 8284. (42) Komarynsky, M. A.; Wahl, A. C. J. Phys. Chem. 1975, 79,695. (43) Haran, N.; Luz, 2.;Shporer, M. J. Am. Chem. SOC.1974,96,4788. (44) Grampp, G.;Jaenicke, W. J. Chem. SOC.,Faraday Trans. 2 1985, 81,1035. (45) Grampp, G.; Jaenicke, W. Ber. Bunsen-Ges. Phys. Chem. 1984, 88, 325. (46) Fuhlendorff, R. Tetrahedron Lett. 1987,28, 5335. (47) Nielson, R. M.; McManis, G. E.; Safford, L. K.; Weaver, M. J. J . Phys. Chem. 1989,93, 2152. (48) Gennett, T.; Milner, D. F.; Weaver, M. J. J. Phys. Chem. 1985, 89, 2787. (49) Weaver, M. J.; Phelps, D. K.; Nielson, R. M.; Golovin, M. N.; McManis, G. E. J. Phys. Chem. 1990,94,2949. (50) Yang, E. S.; Chan, M.-S.; Wahl, A. C. J . Phys. Chem. 1975, 79, 2049. (51) Baranski, A. S.; Winkler,K.;Fawcett, W. R. J.Electroanal. Chem. 1991,313,367.
(52)Bond, A. M.; Henderson, E. L. E.; Mann, D. R.; Mann, T. F.; Thormann, W.; Zoeki, C. G. Anal. Chem. 1988,60,1878. (53) Wipf, D. 0.; Kristensen, E. W.; Deakin, M. R.; Wright”, R. M. Anal. Chem. 1988,60,306. (54) Montenegro, M. I.; Pletcher, D. J . Electroanal. Chem. 1986,200, 371. (55) Nielson, R. M.; McManis, G. E.; Weaver, M. J. J . Phys. Chem. 1989,93,4703. (56) Brunechwig, B. S.; Creutz, C.; Macartney, C. H.; Sham, T. K.; Sutin. N. Faradav Discuss. Chem. SOC. 1982. 74.113. (57) Brown, G.”M.; Sutin, N. J. Am. Chek. Soc. 1979,101, 883. (58) Meyer, T. J.; Taube, H. Inorg. Chem. 1968, 7, 2369. (59) Hupp, J. T.; Weaver, M. J. J. Phys. Chem. 1985,89, 2795. (60) Shporer, M.; Ron, G.; Lowenstein, A.; Navon, G. Inorg. Chem. 1965,4, 361. (61) Wahl, A. C. Z . Elektrochem. 1960,64, 90.
Langmuir, Vol. 9, No. 12, 1993 3615
at defect sites, although one could argue that size selection occurs but is compensated for by a particularly strong adsorption of DMFc. It is significant that even though microscopic defects certainly form at extreme potentials (as indicated by potential-dependent capacitance data), the defects are not so severe as to permit indiscriminate access of all redox-active probes to the electrode surface. For example, the TMPD2+/+reaction appears to be much slower than either the TMPD+/O reaction or any of the ferrocene reactions in spite of the fact that the TMPD2+/+ reaction occurs at a potential outside of the capacitance plateau discussed above. Also, the Fe(CN)63-/k and R u ( N H ~ ) ~ ~reactions + / ~ + involved excursions to potentials outside of the plateau region, yet the voltammetry remained well-behaved and compares quite favorably with that obtained by other workers at similarly coated electrodes studied in water.6**J0 Our generalization from above that interfacial redox reactions of charged species are blocked more effectively than those of neutral species therefore appears to hold even at potentials slightly outside of the capacitance plateau region. This is not to say however that the alkanethiolate monolayers act as ideal barriers with respect to electron transfer at all potentials. For example, the unexpected asymmetry in several of the ferrocene voltammograms (i.e., anodic and cathodic peak potentials are unevenly spaced from the thermodynamic redox potential, see Figures 1and 5) could be accounted for in terms of a potential-dependent partial disruption of the monolayer blocking capacity at positive potentials. Hence, some caution is warranted regarding detailed interpretation of voltammograms occurring at potentials outside of the capacitance plateau region. Summary and Conclusions In summary, we have shown that stable alkanethiolate monolayers can be formed and maintained on gold in propylene carbonate and that, within certain limitations, the alkanethiolate layers possess excellent barrier properties with respect to preventing approach of electroactive solutes to the underlying gold electrode surface. Electrochemical studies on a range of redox-active probe molecules reveal that for a series of ferrocene derivatives the rate of interfacial electron transfer correlates fairly well with the rate of homogeneous electron self-exchange. There is no correlation with molecular size, suggesting that size exclusion at monolayer defect sites is not a major factor in determining electrochemical reactivity at alkanethiolate-coated electrodes under these conditions.The correlation of homogeneous electron self-exchange rates with the observed electrochemical response is poor when structurally different redox systems are compared. It appears that electrically neutral probe molecules generally exhibit more facile interfacial redox kinetics at monolayercoated electrodes than do charged probe molecules. Different degrees of adsorption and/or distances of closest approach for the various redox-activeprobes are postulated to be responsible for this effect. Acknowledgment. Financial support of this work from the National Science Foundation is gratefully acknowledged. K.A.G. acknowledges support from a National Needs Fellowship from the US.Department of Education. Both authors gratefully acknowledgethe skilled assistance of Dr. Chung-Ping Yu in acquiring GC-MS data.
