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Redox Behavior and Ion-Pairing Thermodynamics of Ferrocene and Its Derivatives in the Organic Phase Debo Xiang,†,‡ Guoyu Gao,†,§ Huibo Shao,‡ Hulin Li,§ Hao-Li Zhang,§ and Hua-Zhong Yu*,† Department of Chemistry, Simon Fraser UniVersity, Burnaby, British Columbia V5A 1S6, Canada, Department of Chemistry, Faculty of Science, Beijing Institute of Technology, Beijing 100081, People’s Republic of China, and State Key Laboratory of Applied Organic Chemistry and College of Chemistry & Chemical Engineering, Lanzhou UniVersity, Gansu 730000, People’s Republic of China ReceiVed: September 28, 2009; ReVised Manuscript ReceiVed: NoVember 14, 2009
The redox behavior and ion-pairing thermodynamics of ferrocene and its derivatives (dimethylferrocene and decamethylferrocene) in a nitrobenzene (NB) thin film imposed between a graphite electrode and an aqueous electrolyte solution have been studied. We have shown that the presence of supporting electrolytes in the organic phase complicates their redox behavior, i.e., both the oxidation potential and peak area change significantly with time. In the absence of supporting electrolytes in the NB phase, the redox potential of the ferrocene molecules also shifts when the concentration of the supporting electrolytes (NaClO4 or HClO4) in the aqueous phase varies. It has been confirmed upon considering the influence of liquid-liquid junction potential that such a potential shift is dictated by the formation of ion pairs between the electrochemically generated ferricenium cations and the counteranions in the organic phase. On the basis of the determination of anion (ClO4-) concentrations in the NB film, we were able to quantitatively evaluate the correlation between molecular structure and the formation constant of Fc+ · ClO4- ion pairs. 1. Introduction Ferrocene has been indispensible in contemporary research since its discovery in the 1950s,1 largely due to its inherent organic/inorganic nature including high thermal stability, good solubility in organic media, and reversible redox property. Many derivatives of ferrocene have been synthesized and characterized ever since,2 which accommodated diverse applications including the preparation of functional biomaterials and electronic devices.3-5 Particularly, its favorable redox properties make ferrocene a popular labeling reagent (“mediator”) in the fabrication of electrochemical biosensors; in the past two decades, biosensors based on measuring the electrochemical response of ferrocene-labeled amino acids, peptides, and nucleic acids have attracted much attention.6-10 The oxidation of neutral ferrocene (Fc) results in the formation of ferricenium (Fc+) monocations. Therefore, the redox behavior of ferrocene is inevitably influenced by the type and concentration of anions present in the electrolyte solution.11,12 The ability to form ion pairs with ferricenium cations varies remarkably for different anions, the voltammetric responses of ferrocene (particularly the formal potential and peak shapes) would be changed consequently. There are only a few reports in the literature that provide further insights into this fundamentally important aspect (ion-pairing thermodynamics) for the redox behavior of ferrocene and its derivatives.13-16 The effect of the nature of counterions in aqueous media on the redox behavior of ferrocene was first studied electrochemically in a polymer film by Inzelt and Szabo.13 The ionpair effect on the ferrocene redox reactions was later investigated in self-assembled monolayers (SAMs) that have * To whom correspondence should be addressed. E-mail:
[email protected]. † Simon Fraser University. ‡ Beijing Institute of Technology. § Lanzhou University.
been immobilized on gold electrode surfaces.14-16 Uosaki et al. have shown that the shape and position of the redox peaks of ferrocenyl SAMs are strongly affected by the nature and concentration of the supporting electrolyte, especially the anions in solution.14 Rowe and Creager described the incorporation of ferrocenylhexanethiols into mixed SAMs with different n-alkanethiols (1-CnH2n+1SH, n ) 4, 6, 8, 10, and 12) and their redox behavior in perchloric acid. The propensity of anions to form ion-pairs with ferricenium cations was found to increase upon increasing the chain length of the n-alkanethiols.15 In a later report, Kondo et al. studied the ion-pair formation between ferricenium ions and different anions in organic solvents and discovered a trend opposite to that in an aqueous environment.16 Recently, we have been interested in the study of multicenter ferrocenyl complexes using the “thin-layer electrochemistry” approach, a method developed by Shi and Anson a decade ago to confine redox molecules in an organic film imposed between an aqueous solution and a hydrophobic graphite electrode.17 As of the elimination of diffusion control, we were able to provide further insights into electronic interactions among the redox centers in the same multiferrocenyl molecule18 and to evaluate the correlation between the electron transfer rate across a liquid/ liquid interface and the overall driving force.19 Herein, the thinlayer method was employed to investigate the ion-pair effect on the redox behavior of three structurally different ferrocenyl compounds in the organic phase. In particular, the ion-pair formation constants for these ferrocene derivatives were determined respectively and compared with each other. 2. Experimental Section 2.1. Chemicals. Ferrocene (Fc) was obtained from Acros; dimethylferrocene (DiMFc) and decamethylferrocene (DMFc) were purchased from Aldrich. Tetraethylammonium perchlorate (TEAP) was from Eastman Kodak Co., tetramethylammonium
10.1021/jp909326n 2010 American Chemical Society Published on Web 12/15/2009
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hexafluorophosphate (TMAFP) from Aldrich; hydrochloric acid from Fisher Scientific, sodium hydroxide from Fluka. Sodium perchlorate (98+%) was purchased from Caledon Laboratories Ltd. Nitrobenzene (NB) was obtained from Allied Chemical Ltd.; it was washed consecutively with 0.1 M HCl, 0.1 M NaOH, and deionized water and then saturated with deionized water for at least 12 h before use. All aqueous solutions used for the electrochemical measurements were prepared from deionized water (>18.3 MΩ · cm) previously saturated with NB. 2.2. Apparatus and Procedure. Electrochemical measurements were performed in a conventional three-electrode glass cell. The working electrode is a cylindrical edge-plane graphite (EPG) rod sealed with heat-shrinkable tubing: 0.17 cm2 at the edge of the graphitic planes were exposed. A platinum wire was used as the counter electrode. An Ag|AgCl|3 M NaCl electrode was used as reference in aqueous solution for the thinlayer experiments, while an Ag|Ag+ (0.01 M AgNO3 in NB, saturated TMAFP) electrode was used in NB for the conventional measurements. Before each measurement the working electrode was polished with 400-grit sand paper, washed, and sonicated in deionized water and carefully dried with a heat gun. Then 0.7-1.0 µL of the NB solution of ferrocene derivative was applied to the surface, which spread over the hydrophobic surface quickly to form a thin film. All the electrodes were carefully placed in the aqueous solution. Cyclic voltammetry (CV) was performed with an Autolab Electrochemical Analyzer (PGSTAT30, Eco Chemie BV, Netherlands) in a Faraday cage. All measurements were performed at room temperature. The quantity of HClO4 that partitioned into NB from the aqueous solution was determined by acid-base titrations: mix 35.00 mL of purified NB and an equal volume of HClO4 solution, shake them at least four times and keep them still for 20 h, remove 10 mL NB from the bottom layer into a flask containing 20 mL DMSO, and titrate the HClO4 in the mixture with 1.00 mM fresh Na2CO3 solution. Two drops of bromothymol blue were added as an indicator. The quantity of NaClO4 (concentration of Na+) that partitioned into NB from the aqueous solution was obtained by atomic emission spectrometry. To do this, a series of standard solutions (containing 1.0 to 10 ppm Na+) were prepared for the construction of a calibration curve. 3. Results and Discussion 3.1. Effect of Supporting Electrolytes in the Organic Phase. The thin-layer technique developed a decade ago by Shi and Anson provides a convenient method for studying redox properties of organic compounds and monitoring liquid-liquid interfacial electron transfer processes.17 In particular, examining the CV behavior of redox couples contained in the organic phase allows the determination of the total amount of redox species by utilizing relatively slow scan rates (e.g., 5 mV/s). Since the volume ratio of organic phase to aqueous phase in these experiments is often less than 1/∼104, partitioning of the organic soluble redox couples into the aqueous phase can be troublesome. The other related question is whether it is needed to add supporting electrolytes into the organic phase. We started our investigation by comparing the CV responses of ferrocenyl compounds confined in an NB film on a graphite electrode with and without added supporting electrolytes in the organic phase. Figure 1a shows the CV responses of an EPG electrode covered with 0.9 µL of NB containing 0.80 mM DMFc, for which 0.10 M NaClO4 was employed as the supporting electrolyte in the aqueous solution. In this case, there is no supporting electrolyte in the NB layer. The formal potential does not shift as a function of time; however, the peak area decreases significantly in the
Xiang et al.
Figure 1. CVs of an EPG electrode covered with 0.9 µL of NB (a) containing 0.80 mM DMFc recorded at 0, 3.0, 6.2, and 10.0 min (from outermost to the inside curves) and (b) containing 0.80 mM DMFc and 0.1 M TEAP, recorded at 0, 2.8, 5.8, 9.4, 13.8, and 16.8 min (from outermost to the inside curves). The scan rate was kept as 10 mV/s, and 0.1 M NaClO4 was used as the supporting electrolyte in the aqueous phases for both parts a and b.
first 12 min (from outermost curve to the inside curves). Such a phenomenon has been noted previously in both our report18 and publications from other research groups.20 On the basis of the integration of anodic peaks (Figure 1a), as high as 20% decrease of the DMFc+/DMFc in the NB layer was observed. Figure 1b shows the CV curves obtained with an EPG electrode coated with 0.9 µL of NB containing 0.80 mM DMFc and 0.1 M TEAP, for which 0.1 M NaClO4 was also used as the supporting electrolyte in the aqueous phase. It is clear that the redox peak areas decrease as a function of the time passed since immersion of the electrode in the aqueous solution. In comparison with the results shown in Figure 1a, the change of the integrated charge is significant: within 10 min, it decreased to ∼40% of the initial value. In both cases (parts a and b of Figure 1), the CV response became stable upon establishing ion-partition equilibrium at the interface.20 The remarkable difference between the results shown in parts a and b of Figure 1 is the time-dependent potential shift when supporting electrolytes are added to the organic phase. Initially, the formal potential of the DMFc+/DMFc couple (calculated from the midpoint potential of each CV curve) shifts negatively and becomes constant after 10 min. The decrease of the redox peak areas has been ascribed to the loss of the redox molecules (particularly the ferricenium cations) upon the potential scan, which can be minimized by using hydrophobic ferrocenyl compounds and choosing proper supporting electrolytes.21,22 The time-dependent potential shifts have not been noted previously, which may be due to a change of the liquid-liquid interface junction potential21 and to the ion-pair effect of the ferrocene redox process.11-16 In both cases, a rapid change of the electrolyte concentration in the organic phase in contact with the aqueous phase is expected due to the very small volume of the organic phase. Although the solubility of TEAP in NB is much higher than that in aqueous solution, upon partitioning equilibration its concentration in NB will inevitably drop from
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the original 0.1 M. This in turn leads to a significant decrease of the anion (ClO4-) concentration in the organic phase. The effect of ion-pairing between ferricenium cations (Fc+) and counteranions (ClO4-) plays an important role in the redox properties of ferrocenyl SAMs.14-16 Different concentrations of supporting electrolytes lead to significant variations of the observed formal potentials. Although such an effect has not been discussed in the thin-layer setting, we believe that the ion-pair effect also contributes to the potential shifts shown in Figure 1b.17-19,21-25 Instead of the simple one-electron reversible process (eq 1), the redox reaction should be written more precisely as eq 214-16
Fc - e h Fc+ E°′ Fc +
ClO4 +
K ) [Fc
+
h Fc
· ClO4
(1) -
+e
+ · ClO4 ]/[Fc ][ClO4 ]
(2) (3)
where E°′ is the formal potential of ferricinium/ferrocene redox couple in NB and K is the formation constant of Fc+ · ClO4ion-pairs (as defined in eq 3). On the basis of eq 2, it is clear that the concentration of anions in the organic phase (depending on the partition equilibrium with the aqueous phase) would affect the redox behavior of ferrocenyl compounds; however, the existence of liquid-liquid junction potentials complicates the situation.25 To quantitatively examine the ion-pair effect of ferrocene redox reactions in the organic phase, we will first evaluate the variation of liquid-liquid junction potentials upon changing the concentration of supporting electrolyte in the aqueous phase. 3.2. Liquid Junction Potential at the NB/Water Interface. To study the ion-pairing effect of the ferrocene redox reaction in the organic phase quantitatively (i.e., to determine the ionpair formation constants), we had to systemically adjust the concentration of the anions. This can be achieved by partition equilibration with an aqueous solution of the supporting electrolyte at concentrations of choice. We mixed 35 mL of pure (dry) NB with equal volumes of aqueous NaClO4 and HClO4 solutions at different concentrations (from 0.05 to 2.0 M); the concentrations of NaClO4 and HClO4 in the NB phase upon reaching equilibrium were determined by atomic emission spectrometry and acid-base titration experiments, respectively. As shown in Figure 2, the concentrations of [Na+] and [H+] in the NB phase increase monotonically when the concentrations of NaClO4 and HClO4 in the corresponding aqueous solutions increase from 0.05 to 2.0 M. We assume that NaClO4 and HClO4 are completely dissociated in NB due to their low concentrations (0.02 to 1.0 mM) and the appreciable Ka values (∼3). 25 It has been proposed that the liquid junction potential at an interface is uniquely determined by the initial concentrations and the standard ion transfer potentials when the volume ratio of two phases is extremely small or large.26 With a typical NB/water volume ratio of 1/7.5 × 103, the potential difference of each ion at the interface can be obtained according to eq 427-29
∆wo φ ) ∆wo φ°i + RT/ziF ln (aio/aiw) ≈ ∆wo φi◦ + RT/ziF ln (cio/ciw) (4) where ∆owφ is the Galvani potential difference between the organic (o) and aqueous (w) phases, cio and ciw are the concentrations in the two phases, respectively, and ∆owφ°i is the standard ion transfer potential of the ionic species i, which
Figure 2. (a) Correlation between [Na+] in NB and the concentration of NaClO4 in the adjacent aqueous phase. (b) Correlation between [H+] in NB and the concentration of HClO4 in the adjacent aqueous solution.
TABLE 1: Variation of the Liquid Junction Potential at the NB/Water Interface and the Mid-Point Potential of the DiMFc+/DiMFc Redox Reaction upon Increasing the Concentration of Supporting Electrolyte (NaClO4 or HClO4) in the Aqueous Phase from 0.05 to 2.0 M [ClO4 ]H2O/M
Ej(NaClO4)/V
Ej(HClO4)/V
E1/2(NaClO4)/V vs NHE
E1/2(HClO4)/V vs NHE
0.05 0.1 0.5 1.0 2.0
0.329 0.332 0.349 0.356 0.350
0.318 0.322 0.338 0.338 0.328
0.396 0.384 0.347 0.338 0.326
0.390 0.377 0.341 0.324 0.305
is in fact the standard Gibbs energy of ion transfer (∆Gtr,i°,wfo) expressed on a potential scale28 ◦,wfo ∆wo φi◦ ) ∆Gtr,i /ziF
(5)
By assumption that [Na+]NB ) [ClO4-]NB and that [H+]NB ) [ClO4-]NB and on the basis of the data presented in Figure 2, the specific potential difference of each ion at the interface can be obtained according to eq (4) using the tabulated values of ∆wNBφ°Na+ ) 0.354 V, ∆wNBφ°ClO4- ) -0.083 V, and ∆wNBφ°H+ ) 0.337 V.26,30 The liquid junction potential at the NB/water interface can be readily determined from eq 626 ◦ Ej ) (m∆wo φM + n∆wo φA◦ )/(m + n) m+
n-
(6)
where M A is the only single electrolyte in the two-phase system, Mm+ is the cation, An- is the anion, m and n are their ionic charges, respectively, and ∆wo φ°M and ∆wo φ°A are the potential differences of the cation and anion at the liquid-liquid interface, respectively. As shown in Table 1, the liquid junction potentials for the NB/water system change only slightly ( K(DiMFc) > K(DMFc) in contact with either NaClO4 or HClO4 aqueous solution. Kondo and co-workers have studied the anion effect on the redox behavior of ferrocene-tethered SAMs on gold electrodes in both aqueous and organic solvents (dichloromethane). They proposed that the potential shift was determined by the desolvation energy of the anion (∆G°solv) and the formation energy of the ion-pair (∆G°form).16 The ∆G°solv in an organic solvent is much smaller than that in aqueous solution, so that its contribution to the potential shifts observed in our experiments should be negligible. This means that the ∆G°form dictates the ion-pairing-induced potential shifts in the organic phase. Although ∆G°form includes both electrostatic and van der Waals interactions, the former are typically much stronger, i.e., we need only consider the electrostatic interaction (WE) to explain the ion-pairing effect16
WE ) qFc+qClO-4 /2πεε0d
(9)
where qFc and q are the charges of cation (+1) and anion (-1), respectively, ε0 and ε are the dielectric constants of the vacuum and NB, respectively, and d is the distance between the Fc+ cation and the perchlorate ion in the formed ion pair. It is evident that with the same anion the ion-pair formation constant will depend on the size of ferricenium cation. On the basis of the number of methyl groups on the three ferrocenyl compounds being studied, their radius should decrease in the order rDMFc > rDiMFc > rFc. In this case, the distance between the cation and the perchlorate anion in the ion-pair should increase in the order dFc < dDiMFc < dDMFc. According to eq 9, the interaction between the cation and perchlorate ion decreases in the order WFc > WDiMFc > WDMFc, such a change can well explain the decrease of the ion-pair formation constant (KFc > KDiMFc > KDMFc) observed above. The ion-pair formation constants for Fc and DiMFc in NaClO4 are smaller than those in HClO4 while the value for DMFc in NaClO4 is slightly larger than that in HClO4. We are not sure at this stage about the origin of such a difference, which may relate to the fact that the dissociation of HClO4 in NB is less effective than that of NaClO4. It is interesting to note that the ion-pair formation constants obtained herein are much larger than those obtained by Rowe and Creager for mixed ferrocenylhexanethiol/n-alkanethiol SAMs.15 They have shown that the formation constants increase by a factor of 25 for the coadsorbed n-alkanethiols that are from 6 to 12 carbons long. It has been suggested that as the environment surrounding the ferrocene moiety becomes more alkanelike with increasing alkanethiol coadsorbate chain length, ion pairing of ferricinium with perchlorate ion becomes more favorable. The even larger formation constants we have obtained in the organic phase provided further evidence for the above hypothesis. It should be also noted that their “effective” formation constants have been obtained in the presence of 1.0 M H2SO4.15 The thinlayer experiments enabled us to evaluate the ion-pairing effect in the presence of a single type of anions. +
ClO4
4. Conclusion The presence of a supporting electrolyte in the organic phase complicates the redox properties of ferrocenyl compounds in thin-layer experiments; the partition equilibration with the aqueous phase containing supporting electrolytes is sufficient to obtain stable and reproducible CV responses. The potential shifts observed for the ferrocenyl compounds in the organic
phase upon increasing either HClO4 or NaClO4 concentrations in the aqueous solution are dictated by the Fc+ · ClO4- ionpairing thermodynamics and influenced by the variation of the liquid junction potential. The calculated formation constants are in the order of 106 to 109, indicating that stable Fc+ · ClO4pairs are formed in the organic phase. In addition, the ionpair formation constant decreases in the order of KFc > KDiMFc > KDMFc, which can be explained by the expected differences in the electrostatic interaction between Fc+ and perchlorate ion in NB. Acknowledgment. The authors gratefully acknowledge the financial support from the Natural Sciences and Engineering Research Council of Canada. D.X. and G.G. thank the State Scholarship Fund of China to support their stay as visiting graduate students at Simon Fraser University in Canada. H.Z.Y. thanks Dr. Eberhard Kiehlmann for proofreading the revised manuscript. References and Notes (1) Kealy, T. J.; Pauson, P. L. Nature 1951, 168, 1039–1040. (2) Ferrocenes: Ligands, Materials and Biomolecules; Stepnicka, P., Eds.; John Wiley & Sons, Inc.: New York, 2008. (3) Tour, J. M. Acc. Chem. Res. 2000, 33, 791–804. (4) Ward, M. D. Chem. Soc. ReV. 1995, 24, 121–134. (5) Astruc, D. Electron transfer and radical processes in transitionmetal chemistry; Wiley-VCH; Marcel Dekker: New York, 1995. (6) Takenaka, S.; Uto, Y.; Kondo, H.; Ihara, T.; Takagi, M. Anal. Biochem. 1994, 218, 436–443. (7) Ihara, T.; Maruo, Y.; Takenaka, S.; Takagi, M. Nucleic Acids Res. 1996, 24, 4273–4280. (8) Uto, Y.; Kondo, H.; Abe, M.; Suzuki, T.; Takenaka, S. Anal. Biochem. 1997, 250, 122–124. (9) Takenaka, S.; Yamashita, K.; Takagi, M.; Uto, Y.; Kondo, H. Anal. Chem. 2000, 72, 1334–1341. (10) Kraatz, H.-B. J. Inorg. Organomet. Polymer Mater. 2005, 15, 83– 106. (11) Yang, E. S.; Chan, M.-S.; Wahl, A. C. J. Phys. Chem. 1980, 84, 3094–3099. (12) Hupp, J. T. Inorg. Chem. 1990, 29, 5010–5012. (13) Inzelt, G.; Szabo, L. Electrochim. Acta 1986, 31, 1381–1387. (14) Uosaki, K.; Sato, Y.; Kita, H. Langmiur 1991, 7, 1510–1514. (15) Rowe, G. K.; Creager, S. E. Langmiur 1991, 7, 2307–2312. (16) Kondo, T.; Okamura, M.; Uosaki, K. J. Organomet. Chem. 2001, 637, 841–844. (17) Shi, C.; Anson, F. C. Anal. Chem. 1998, 70, 3114–3118. (18) Wang, M. C. P.; Li, Y.; Merbouh, N.; Yu, H.-Z. Electrochimi. Acta 2008, 53, 7720–7725. (19) Xu, J.; Frcic, A.; Clyburne, J. A. C.; Gossage, R. A.; Yu, H.-Z. J. Phys. Chem. B 2004, 108, 5742–5746. (20) Shafer, H. O.; Derback, T. L.; Koval, C. A. J. Phys. Chem. B 2000, 104, 1025–1032. (21) Shi, C.; Anson, F. C. J. Phys. Chem. B 1999, 103, 6283–6289. (22) Shi, C.; Anson, F. C. J. Phys. Chem. B 2001, 105, 8963–8969. (23) Shi, C.; Anson, F. C. J. Phys. Chem. B 1998, 102, 9850–9854. (24) Shi, C.; Anson, F. C. J. Phys. Chem. B 2001, 105, 1047–1049. (25) Chung, T. D.; Anson, F. C. Anal. Chem. 2001, 73, 337–342. (26) Kakiuchi, T. Anal. Chem. 1996, 68, 3658–3664. (27) Hornolka, D.; Hung, L. Q.; Hofrnanova, A.; Khalil, M. W.; Koryta, J.; Marecek, V.; Samec, Z.; Sen, S. K.; Vany´sek, P.; Weber, J.; Brezina, M.; Janda, M.; Stibor, I. Anal. Chem. 1980, 52, 1606–1610. (28) Reymond, F.; Steyaert, G.; Carrupt, P.-A.; Testa, B.; Girault, H. J. Am. Chem. Soc. 1996, 118, 11951–11957. (29) Quinn, B. M.; Ding, Z.; Moulton, R.; Bard, A. J. Langmuir 2002, 18, 1734–1742. (30) Vany´sek, P. L.; Ramı´rez, B. J. Chil. Chem. Soc. 2008, 53, 1455– 1463. (31) Long, H. C. D.; Buttry, D. A. Langmuir 1990, 6, 1319–1322. (32) Ball, J. C.; Marken, F.; Wadhawan, J. D.; Fulian, Q.; Wadhawan, J. D.; Blythe, A. N.; Schro¨der, U.; Compton, R. G.; Bull, S. D.; Davies, S. D. Electroanalysis 2000, 12, 1017–1025. (33) Rayner, D.; Fietkau, N.; Streeter, I.; Marken, F.; Buckley, B. R.; Page, P. C. B.; Campo, J. d.; Mas, R.; Muoz, F. X.; Compton, R. G. J. Phys. Chem. C 2007, 111, 9992–10002.
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