J. Phys. Chem. 1994,98, 3843-3850
3843
Kinetics of Electron Transfer to Attached Redox Centers on Gold Electrodes in Nonaqueous Electrolytes Melissa S. Ravenscroft and Harry 0. Finklea' Department of Chemistry, West Virginia University, Morgantown, West Virginia 26506 Received: October 21, 1993; In Final Form: February 1 , 1994'
Self-assembled monolayers with attached redox centers are characterized by their reversible behavior and their electron-transfer kinetics in a range of nonaqueous solvents: methanol, ethanol, propanol, butanol, acetonitrile, N,N-dimethylformamide, dimethylsulfoxide, propylene carbonate, acetone, and tetrahydrofuran. The monolayers are formed by coadsorbing the thiols HS(CHZ),CONHCHZ~~R~(NH~)~(PF& and HS(CHz),COOH ( n = 10 or 15) on gold electrodes. The monolayer-coated electrodes are examined by cyclic voltammetry and chronoamperometry first in the nonaqueous electrolyte and then in an aqueous electrolyte. The reversible CV's (n = 10) indicate the presence of strong ion-pairing and the relatively disordered structure of the monolayers in the nonaqueous solvents. A method is introduced for the correction of i R drop distortion when the chronoamperometry data are analyzed for rate constants. Apparent rate constants (n = 15) are obtained as a function of the percent conversion of the redox centers from the initial to the final oxidation state and of the iR drop-corrected overpotential. Apparent standard rate constants and reorganization energies are obtained by fitting the Tafel plots with Marcus theory. For most of the more polar solvents, the apparent reorganization energies are nearly identical to the values obtained in water (0.9 eV for oxidation and 0.7-0.8 eV for reduction of the redox centers); the least polar solvents yield lower apparent reorganization energies. Propylene carbonate data deviate markedly from the pattern of the other polar solvents. The standard rate constants in water are reproducibly close to 1 s-l. The apparent standard rate constants in the nonaqueous solvents show a considerably greater heterogeneity and are generally faster by up to a factor of 10 than the corresponding aqueous standard rate constants; however, the standard rate constants do not correlate with solvent relaxation times. The dominant factors which control the kinetic parameters of the monolayers in nonaqueous solvents appear to be monolayer disorder and the local water concentration.
Introduction Redox centers which are incorporated into a self-assembled thiolate monolayer (i.e. metal/monolayer/redox centers) provide a probe of both monolayer structure and the kinetic parameters of long-range electron transfer.l-I6 Generally the redox centers are diluted by forming a mixed monolayer system composed of both the electroactive thiol and a diluent thiol in order to inhibit electron transfer between adjacent centers. Most commonly the redox centers are ferrocenes, viologens, or quinones. We'' have found that pyRu(NH3)52+/3+(py = pyridine) is an excellent redox center for the study of electron-transfer kinetics as a function of overpotential, length of the hydrocarbon chain, electrolyte composition, and temperature. Electrochemical methods (cyclic voltammetry and chronoamperometry) provide information on both the order of the monolayer and the electron-transfer kineticsbetween the electrode and the redoxcenter. Cyclicvoltammetry (CV) yields the formal potential ( E O ' ) and the coverage (Q) of the redox centers in the monolayer. Under reversible conditions, a cyclic voltammogram with zero peak splitting (Up) and 90-100 mV peak half-width (AEfwhm) indicates a homogeneous environment around the redox centers (the formal potentials of all of the redox centers are the same), and hence an ordered monolayer can be inferred. Chronoamperometry(CA) data obtained for a range of potential steps yield rate constants for electron transfer which can be used to construct a Tafel plot (In k vs 7,where 7 = E - E O ' ) for a given monolayer. The Tafel plot can then be fitted with the Marcus theory of electron transfer to yield the standard rate constant and the reorganization energy (see below). The behavior of mixed monolayers composed of HS(CH~),CONHCH~~~RU(NH~)~~+/~+ (abbreviated as HS*Abstract published in Advance ACS Abstracts, March 15, 1994.
0022-3654/94/2098-3843%04.50/0
C,-Ru) and HS(CHz),,COOH (abbreviated as HS-C,,-COOH), where n = 10, 11, or 15, has been extensively studied in aqueous electrolyte^.'^ In particular, the HS-CI~-RU/HS-C~~-COOH monolayer system is stable and well-behaved in weakly acidic aqueous electrolytesup to 60 "C. The results are consistent with a well-oriented monolayer with thepyR~(NH3)5~+/~ redox centers residing in the aqueous phase. Cyclic voltammetry, chronoamperometry, and AC impedance spectroscopy have all been used to establish and verify a standard rate constant of ca. 1 s-1 for the HS-C 5-Ru/ HS-C I 5-COOH system. Nonaqueous solvents provide a number of useful tests for monolayer structure and electron-transfer theory because of their range of properties (liquidus temperatures, dielectric constants, polarity, relaxation times). The highly ordered structure of alkanethiolatemonolayers in water may be lost as less polar solvent molecules interact with the alkane chains. The solvent properties can substantially affect electron-transfer kinetics. The following discussion outlines the theory used to fit experimental Tafel plots and defines the specific solvent influences. The equations are limited to the case of reduction of oxidized redox centers. The cathodic rate constant can be predicted by an integration of donor and acceptor levels over a range of energies E (defined with respect to the Fermi energy of the metal):2bJ7
The donor levels are the occupied energy levels of the metal as given by the Fermi function n(E) (eq 3); the density of states 0 1994 American Chemical Society
3844
Ravenscroft and Finklea
The Journal of Physical Chemistry, Vol. 98, No. 14, 19‘94
in the metal is assumed to be constant over the range of energies needed to evaluate the integral. The oxidized redox centers are modeled by a Gaussian distribution of electron acceptor levelsD,,(E) whose position with respect to the formal potential and whose width are defined by the reorganization energy X (eq 2). The reorganization energy is further divided into an inner-sphere component Ab, dependent only on the identity of the redox center, and an outer-sphere component A,. A frequently used expression1*for X, is
where e is the electron charge, NA is Avogadro’s constant, a is the radius of the redox center, d is the distance between the redox center and electrode, copis the optical frequency dielectric constant (usually the square of the index of refraction for visible light), and c, is the static dielectric constant of the solvent. Hence, the rate of electron transfer should be sensitive to the solvent for redox centers in which the outer-sphere component of the reorganization energy dominates. Some successhas been achieved in finding correlations of rate constants with solvents predicted by this relationship in homogeneous systems,18especially in mixedvalence complexes.19 A tunneling factor P(E) is included in the integral. As shown previously,17ait decreases exponentiallywith distance d (P(E)= exp(-Bd), where 0 is the tunneling constant) between the redox centers and the metal. However, the functional dependence of P(E) on the height of the tunneling barrier or the energy at which electron transfer occurs is not known. The average barrier height does not appear to change with electrode potential, contrary to the prediction of a model based on quantum mechanical tunneling through a dielectric 1 a ~ e r . l ’ ~ The solvent longitudinal relaxation time T L can appear in the pre-integral factor A with an adjustable exponent, Le. TLO with 19 varying from 0 to 1.20,21For nonadiabatic electron transfer (e.g. long-range electron tunneling), this factor is expected to be negligible.21,22 We present a preliminary survey of the effects of nonaqueous solvents on the HS-ClrRu/HS-Cl+200H and the HS-ClsRu/HS-C15-COOH systems. This extension of electroactive self-assembled monolayers into the realm of nonaqueous solvents will address the issueof solvent effects on electron-transferkinetics and monolayer order. Because electrolyte resistance is considerably greater in the nonaqueous solvents, we introduce a new method for analyzing CA data for rate constants which corrects for iR drop. The new analysis changes some of the previously published values for the apparent reorganization energies. Experimental Section The working electrode in all cases was a polycrystalline gold flag. The gold flag was cleaned by heating to incandescence in a gas-air flame. The electrochemical cell had a gold wire counterelectrodein a coil around the gold flag electrode. A 1-mm i.d. Luggin capillary connected the SSCE (NaCl saturated) reference electrode to a point within 0.5 cm of the gold flag. CV’s were acquired using a Bioanalytical Systems CV-27 potentiostat and a Zenith XT computer with a Metrabyte DASH16 ADC board. Current andvoltage signals were analyzed using ASYSTANT software (MacMillan) to extract the electrode coverage (Q), formal potential ( E O ’ ) , peak splitting (Up), and peak width at the full width at half-maximum (AEfwhm). Chronoamperometry experiments were performed using an EG&G Princeton Applied Research Model 273 potentiostat and EG&G Model 270 Electrochemical Analysis Software. HS(CHz),CONHCH2pyRu(NH3)52+(where n = 10 or 15), CH~CONHCH~~~RU(N (the H ~solution ) ~ ~ + analog redox center), and HS(CH2),COOH (wheren = lOor 15) weresynthesized as previously described.I7 The procedure for generating mono-
+
layers with the smallest and most uniform rate was slightly modified. Instead of collecting CV’s in 1 M Na2S04 between depositions, the CV’s were collected in the appropriate nonaqueous solvent between depositions. After monolayer formation, the experimental protocol consisted of the following: (a) CV in nonaqueoussolvent, (b) CA in nonaqueous solvent (stepping the applied overpotential t) symmetrically about E O ’ , with Eo’ + t) as the initial potential step), (c) CV in aqueous Na2S04 (1 M, pH 4), and (d) CA in aqueous Na2S04 (same protocol as in step b). The data collected in the aqueous electrolyte served as a reference for the kinetics in the nonaqueous solvent and as a check for monolayer structure and composition. The formal potentials needed for the the potential program of the CA experimentswere obtained from the immediatelypreceding CV’s. The following 10 nonaqueous solvents were used as received: acetonitrile (AN), acetone (AC), methanol (MeOH), ethanol (EtOH), n-propanol (PrOH), n-butanol (BuOH), N,N-dimethylformamide (DMF), dimethylsulfoxide (DMSO), propylene carbonate (PC), and tetrahydrofuran (THF). For thenonaqueous solvents, the electrolyte used was either 0.1 M LiC104 (AN, AC, MeOH, EtOH, PrOH, BuOH) or 0.1 M TBAP (tetrabutylammonium perchlorate) (DMF, DMSO, THF, PC). TBAP was twice recrystallized from an ethanol/water mixture. No special efforts were made to dry the components of the electrolytes or limit exposure of the electrolyte to ambient air other than normal purging with argon; consequently, water was undoubtedly present at millimolar or greater concentrations. The uncompensated resistance R, and RC cell time constant were determined by analyzing the non-Faradaic current transient for a 0.1-V potential step at potentials positive of the formal potential of the redox centers (n = 15). Approximate R, values in ohms were as follows: H20 (4); AN (14); AC (48); MeOH (56); DMF (74); EtOH (150); DMSO (180); PC (190); PrOH (360); T H F (450); BuOH (750). These resistances were used to correct the CA data for iR drop (see below). In all cases the RC time constants were sufficiently small that the Faradaic current was easily distinguished from the charging current transient. Results and Discussion Cyclic Voltammetry. The stability of the redox centers is a critical problem in nonaqueous solvents. In our initial experiments, the monolayers were formed and characterized in aqueous Na2S04and then examined in the nonaqueous electrolyte. We always observed a marked loss (on the order of 50%) of the integrated charge Q in the first nonaqueous CV relative to the preceding aqueous CV. Since the charging currents were reproducible as the electrode was transferred from the nonaqueous to the aqueous solvent (and viceversa), we inferred that the redox centers were being lost while the alkanethiol layer was being retained. The mechanism for the loss of redox centers is not known. In order to minimize loss in coverage between solvents, it is necessary to characterize each monolayer first in the nonaqueous solvent and then in the aqueous electrolyte. This protocol results in reasonable reproducibility of the charge Q in most cases. Table 1 lists the coverages obtained from the final CV in the nonaqueous solvent and the initial CV in the aqueous solvent. The CA experiments in the nonaqueous solvent are performed between these two measurements. In almost all cases the coverage decreases by no more than 20%, and we are reasonably confident that the kinetic measurements are performed on the same population of redox centers in the two solvents. The major exception is propylene carbonate in which the coverage appears to continually drop while the monolayer is exposed to this solvent; there are severalother anomalous results in propylene carbonate (see below). The composition of the deposition solutions (typically 1:l to 1:3 redox-to-diluent thiol mole ratio) is deliberately adjusted to
Electron Transfer Kinetics
TABLE 1: Stability of Redox Center Coverages in Nonaqueous Electrolytes. Q cathode (pC) Q anode (PC) solvent nonaq aq nonaq aq PC 6.4 3.1 6.3 2.6 DMSO 3.7 2.9 3.4 2.7 AN 3.4 3.1 3.8 2.9 DMF 3.O 3.0 2.9 2.9 MeOH 3.8 3.6 3.9 3.6 EtOH 3.7 3.2 4.0 3.1 AC 3.1 2.9 2.8 2.8 PrOH 3.0 3.0 3.1 2.9 BuOH 5.2 4.9 6.3 4.1 THF 5.3 4.2 5.1 4.0 a CV coverages (Q) are shown for both the anodic and cathodic peaks monolayers. The for each of 10 different HS-C~~-RU/HS-CI~-COOH nonaqueous and aqueous coverages are each determined before the CA experiment in the respective electrolyte. Scan rate = 0.1 V/s.
The Journal of Physical Chemistry, Vol. 98, No. 14, 1994 3845 TABLE 3: Reversible A HS-Clo-Ru/HS-Clo-COO Aqueous Electrolytes' nonaqueous 4 Unvnu solvent (mV)
A E m Values for %andMonolayers in Nonaqueous and aqueous
" w m " v n u &-I an. (mV) cat. (mV) (mV) an. (mV) cat. (mV) (uCI 114 121 2 93 93 2.7 100 99 2 95 94 1.8 116 116 0 102 92 1.1 139 160 2 98 94 1.8 147 149 1 98 95 2.2 159 167 2 101 94 2.0 117 127 4 98 94 1.5 222 204 2 110 98 1.7 247 286 2 100 95 3.4 201 202 4 112 101 1.9 u
p
~~~
AN AC MeOH EtOH PrOH BuOH DMSO DMF THF PC a
2 0
9 36 36 32 22 101 72 119
QTotal is the final coverage in H20. Scan rate = 0.1 V/s.
the perchlorate anion than the Ru" species, the Eo'surf will shift negative because the RuII'species is stabilized relative to the RuII species. Since the nonaqueous E"' surface values are all substantially shifted relative to the Eo' solution values, it is likely EO' EO' AEO' that ion-pairing tendencies are greatly enhanced for the surfacesolution HS-Cls-Ru AEO' HS-ClwRu A E O ' attached redox centers in nonaqueous solvents. One anomaly is solvent (V vs SSCE) (V vs SSCE) (V) (V vs SSCE) (V) the very large negative shift in E"' in propylene carbonate (300-0.020 -0,003 -0,020 -0.003 Hz0 -0.017 400 mV) despite the high static dielectric constant (es = 64) of PC 0.233 -0.209 -0.442 -0.079 -0.312 this solvent. DMSO 0.018 -0.17 -0.19 -0.160 -0.178 AN 0.250 0.17 -0.08 0.174 -0.076 As noted in the Experimental Section, no attempt was made DMF 0.103 -0.17 -0.27 -0.200 -0.303 to rigorously dry the nonaqueous solvents in these experiments. MeOH 0.175 -0.01 -0.18 -0.059 -0.234 The local water concentration at the monolayer/solvent interface EtOH 0.18 -0.02 -0.20 -0.033 -0.208 may be enhanced relative to the solution concentrations because AC 0.21 0.10 -0.11 0.200 -0.009 of the charged and polar character of the monolayer surface. PrOH 0.166 0.04 -0.12 0.021 -0.145 However, the large shifts in formal potential indicate that the BuOH 0.15 0.039 -0.11 0.044 -0.110 THF 0.232 0.01 -0.22 0.048 -0.184 local environment around the redox centers in the nonaqueous solvents is different from that found in a pure aqueous electrolyte. a A E O ' = AEofSurfAEo'lol. The least significantdigit in each formal The peak splitting and the peak half-width of a reversible cyclic potential reflects the reproducibility (less than or greater than 5 mV) of the formal potentials over at least two independent measurements. voltammogram are other measures of monolayer behavior. Reversible behavior requires very slow scan rates for the HSachieve the coverages shown in Table 1. At these coverages the C1s-Ru/HS-C1s-COOH monolayers with concomitant noisier Faradaic current is easily measured relative to the charging current data, so a shorter chain system is used to investigate Upand in the CA experiments while the average spacing between redox UFWHM (Table 3). For a mixed monolayer system consisting centers is large enough (4 pC/cm2 correspond to 20 8,spacing) of HS-ClrRu and HS-Cl&OOH, reversible behavior (Up to avoid possible interference in the kinetic measurements by
