Communication pubs.acs.org/Organometallics
Cyclic Voltammetric Studies of Chlorine-Substituted Diiron Benzenedithiolato Hexacarbonyl Electrocatalysts Inspired by the [FeFe]-Hydrogenase Active Site Elizabeth S. Donovan,† Joseph J. McCormick,† Gary S. Nichol,‡ and Greg A. N. Felton*,† †
Department of Chemistry, Oakland University, 2200 North Squirrel Road, Rochester, Michigan, United States School of Chemistry, The University of Edinburgh, King's Buildings, West Mains Road, Edinburgh, U.K.
‡
S Supporting Information *
ABSTRACT: Chlorine-substituted benzenedithiols have been coordinated to iron carbonyl to yield a series of hydrogen-producing electrocatalysts: [Fe2(μS2C6H4−xClx)(CO)6], where x = 0, 2−4, and [Fe2(μ-S2C6H3Me)(CO)6]. Among this series the novel compounds [Fe2(μ-S2C6Cl3H)(CO)6] (4) and [Fe2(μ-S2C6Cl4)(CO)6] (5) have been characterized by X-ray crystallography. IR spectroscopy and electrochemical analysis were used to evaluate the electronic structure of these proton-reducing molecular electrocatalysts. The increase in the number of electron-withdrawing groups upon the electrocatalyst leads to an anodic shift in reduction potential and a concomitant lowering of the overpotential of hydrogen evolution. An overpotential of just 0.39 V is observed for 5 with acetic acid (pKa = 22.3 in acetonitrile).
D
The use of 3 in multicomponent photocatalytic14 systems and recent similar use of 215 suggest a need for more easily reducible electrocatalysts; therefore, incrementally adding chlorine substituents to the benzenedithiolato bridge of 2 has been studied herein.
evelopment of a viable earth-abundant molecular electrocatalyst1 for producing dihydrogen is of interest, since H2 is a carbon-neutral renewable fuel able to replace fossil fuels.2 Since current dihydrogen production does not meet cost and efficiency needs for development of a “hydrogen economy”,3 efforts are being made to mimic the proton-reducing efficiency of hydrogenase enzymes4 (eq 1). Due to the structural similarity between the active site of [FeFe]-hydrogenase (Figure 1)5 and Reihlen-type6 molecules, electrocatalytic
2H+ + 2e− ⇌ H 2
Benzenedithiols were prepared/purchased and were allowed to reflux for 70−90 min in THF with triiron dodecacarbonyl, which resulted in stable Reihlen-type complexes as red-orange solids (Figure 2), with confirmation of novel compound structures by X-ray crystallography (Figure 3 and Table 1). A comparison of the X-ray determined bond lengths and angles (Table 1) does not display any clear trends; however, there is a strong correlation between IR absorption (Figure 4) of the carbonyl groups of each compound and the reduction potential (Figure SI-3, Supporting Information). This correlation follows the expected trend of electron-withdrawing groups removing electron density from the metal center and a subsequent lessening of the π back-bonding from the metal into the M−CO bond, supporting the assignment of a primarily metal-centered LUMO. The change in reduction potential observed with the increase in the number of electronwithdrawing groups does not strictly follow an additive pattern. Compound 3 (two Cl groups) is easier to reduce by 110 mV than 2 (no Cl groups), and the shift to 4 (three Cl groups) is 50 mV easier to reduce than 3, yet the increase in chlorination
Figure 1. Proposed active-site structure of [FeFe]-hydrogenase.
evaluations of species with the general formula [Fe2(μSR) 2(CO)nL6−n] (n = 0−6) has flourished.7 Although molecules with this formula have been shown to evolve dihydrogen in the presence of certain acids,8 they do so at much slower rates and overpotentials significantly higher than [FeFe]-hydrogenase.9 Use of an aromatic bridging thiolate has been explored,10 with several examples of improved overpotential due to the use of electron-withdrawing groups on the dithiolate bridge.11 Indeed, the reports of electrochemical reversibility and proton reduction by Fe2(μ-S2C6H4)(CO)6 (2)12 led to the synthesis of Fe2(μ-S2C6Cl2H2)(CO)6 (3).13 © 2012 American Chemical Society
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Received: October 5, 2012 Published: November 12, 2012 8067
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Figure 2. Fe2S2-hexacarbonyl compounds [Fe2(μ-1,2-S2C6H3-4-CH3(CO)6]16 (1), [Fe2(μ-1,2-S2C6H4)(CO)6]17 (2), [Fe2(μ-1,2-S2C6H2-3,6Cl2)(CO)6]13 (3), [Fe2(μ-1,2-S2C6H-3,4,6-Cl3(CO)6] (4), and [Fe2(μ-1,2-S2C6Cl4)(CO)6] (5).
Figure 3. ORTEP molecular graphic showing the structures of 4 and 5 with displacement ellipsoids at the 50% probability level.
Table 1. Selected Bond Lengths (Å), Bond Angles (deg), IR Absorptions (cm−1, Recorded in Hexane), and E° (V vs Fc+/ Fc) Values for Compounds 1−5 param
118
217
313
4
5
Fe−Fe Fe1−S2 Fe2−S1 Fe1−S1 Fe2−S2 C−Oa S−Fe−Sa Fe−S−Fea CO
2.475 2.265 2.269 2.262 2.271 1.139 80.4 66.2 2005.8 2044.2 2078.8 −1.34b
2.480 2.271 2.272 2.262 2.267 1.135 80.7 66.3 2006.8 2045.2 2079.7 −1.31c
2.480 2.266 2.267 2.267 2.256 1.128 81.2 66.2 2011.8 2049.3 2083.0 −1.20d
2.476 2.274 2.271 2.268 2.273 1.142 81.2 66.1 2013.2 2050.3 2083.7 −1.15
2.480 2.273 2.267 2.265 2.273 1.134 81.2 66.2 2014.2 2050.9 2084.4 −1.13
E°
Figure 4. IR spectra in hexane of compounds 1 (blue) and 5 (green) highlighting the extremes of the series studied, with an increase in electron-withdrawing substituents (IR spectra in hexane of all compounds 1−5 can be seen in Figures SI-1 and SI-2 in the Supporting Information).
Average value. bOtt et al. reported −1.33 V vs Fc+/Fc.13 cCapon et al. reported −1.27 V vs Fc+/Fc,12a Evans et al. reported −1.32 V vs Fc+/ Fc,12c and Ott et al. reported −1.31 V vs Fc+/Fc.13 dOtt et al. reported −1.20 V vs Fc+/Fc.13
a
comparison to the closer enzymatic structural mimicry of three-atom bridged dithiolates.20 Variation of the CV scan rate21 (Figures SI-4−SI-9, Supporting Information) was probed to see if the effect of decreased electron density upon the iron atoms would lead to an enhanced rate of Fe−S cleavage. A steeper slope of a plot of normalized peak heights against log scan rate (Figure SI-9) is observed for the chlorinated compounds 3−5, with slopes ∼20% steeper than those observed for 2 and ∼60% steeper than those observed for 1. A steeper slope indicates a slower scan rate at which a one-
to give 5 (four Cl groups) only yields a further 20 mV increase in the ease of reduction. Indeed, this effect has been seen with benzenedithiolate (bdt) ligands (attached to CpCo), where Cl2bdt to Cl3bdt sees an anodic shift in reduction potential of 70 mV, while Cl3bdt to Cl4bdt only sees a shift of 20 mV.19 The reversible electrochemistry of compounds 1−5 is highlighted in Figure 5, a reversibility that is thought to be one of the key advantages13 of benzenedithiolate-bridged [FeFe]-hydrogenase active site inspired electrocatalysts in 8068
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Table 2. Observed Overpotential, Turnover Frequency (TOF), and Catalytic Efficiency (CE) for Compounds 1−5 in the Presence of 50 mM Acetic Acid overpotential (V) “TOF” (s−1)b CEb
a
1
2
3
4
5
0.64 23.6 0.44
0.54 25.6 0.46
0.51 12.8 0.32
0.42 11.8 0.31
0.39 6.5 0.23
Calculated using −1.46 V vs Fc+/Fc as the standard potential for reduction of acetic acid.24 bMethodology elaborated in the Supporting Information. a
(pKa = 22.3 in acetonitrile). The catalytic waves with acetic acid (Figures SI-10−SI-14, Supporting Information) for compounds 3−5 are somewhat closer to the theoretical “S-shape” of an unperturbed catalytic response, suggesting less catalyst decomposition/inhibition on the CV time scale in comparison to compounds 1 and 2.25 Turnover frequency (TOF)1 is frequently used as a metric for describing proton-reducing catalysts, commonly based upon the ratio of catalyst peak height to catalytic peak height (using cyclic voltammetry), measured in the acid-independent region, to obtain a pseudo-first-order rate constant. In the case of compounds 1−5 an acid-independent region was not observed, suggesting a component of direct reduction at high values of excess acetic acid (100-fold excess). Low turnover number (TON) also affects the determination of the TOF (the TON for compound 2 is ∼11). Therefore, a “relative TOF” was calculated with 50 mM acetic acid with values of 7−24 s−1 as chlorination is decreased (compounds 5−1). The magnitude of these numbers may be used to compare to other catalytic systems, and the series trend is as expected, given that protonation of the dianion is the rate-limiting step12c (a greater number of electron-withdrawing groups makes the catalyst less basic and subsequently harder to protonate). However, the series trend is also affected by the differing overpotential of catalysis in each case; for example, compound 5 has the lowest relative TOF but also the lowest overpotential driving the reaction. A further form of assaying relative catalytic efficiency (CE)1,7b is presented in Table 2, similarly showing higher efficiency for the unchlorinated compounds (albeit at higher overpotential), pushing CE into the moderate category. The CV of the catalysts in the presence of a stronger acid (ptoluenesulfonic acid, pKa = 8.7 in acetonitrile) was briefly probed: while electrochemical reversibility of the first (noncatalytic) peak is lost, indicating protonation (Figure SI-15, Supporting Information), direct reduction of the acid in this region makes any analysis of catalytic activity questionable (Figure SI-16, Supporting Information). To obtain more accurate TOF values an attempt was made to study the catalysts upon a mercury film working electrode (see the experimental section and Figures SI-17−SI-21 in the Supporting Information). A mercury electrode has no direct reduction of acetic acid at the potentials of the catalytic peak for these catalysts, and therefore such direct reduction would not interfere with the observation of an acid-independent catalysis region. This approach is appropriate for compounds 1 and 2, but the chloro groups of compounds 3−5 interact strongly with the mercury surface. In conclusion, the use of increasing numbers of electronwithdrawing groups upon the dithiolate bridge of [FeFe]hydrogenase-inspired electrocatalysts lowers the overpotential of catalytic reduction of protons from weak acids. Such
Figure 5. CV of compounds 1−5 displaying the anodic potential shift of reduction as electron-withdrawing substituents are increased (scan rate 0.100 V s−1, glassy carbon 3 mm diameter working electrode in 0.100 M Bu4NPF6 anhydrous acetonitrile).
electron process (where Fe−S cleavage has not had time to occur)22 is obtained. Thus, a reduced rate of Fe−S bond breakage for the more chlorinated compounds is observed, suggesting a steric effect overcoming the anticipated weakening of the Fe−S bond due to the more electron-poor iron centers of 3−5. The catalytic cycle for compounds 1−5 is assumed to proceed via the same mechanism reported for catalysis of weak acid by compound 2.12c The catalytic reduction of protons from weak acids can be described to proceed via an E(ECEC) mechanism (shown in the Supporting Information) and has been computationally demonstrated to proceed via second protonation upon an available sulfur (after an Fe−S bond breakage).23 There is an anodic change in the reduction potential of catalysts 1−5 of 210 mV as the number of electronwithdrawing groups is increased. Unsurprisingly, this leads to a similar magnitude 250 mV anodic change in the position of the catalytic peak (Figure 6), and subsequent 250 mV improvement of the overpotential24 of acetic acid reduction (Table 2). An overpotential of just 0.39 V is observed for 5 with acetic acid
Figure 6. CV of compounds 1 and 5 at 1.00 mM in the presence of 12.50 mM acetic acid at 0.100 V s−1 scan rate. The measurement for 12.50 mM acetic acid without catalyst present is shown by a dotted gray line. 8069
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(6) Reihlen, H.; Gruhl, A.; von Hessling, G. Justus Liebigs Ann. Chem. 1929, 472, 269−287. (7) (a) Capon, J.-F.; Gloaguen, F.; Petillon, F. Y.; Schollhammer, P.; Talarmin, J. Coord. Chem. Rev. 2009, 253, 1476−1494. (b) Felton, G. A. N.; Mebi, C. A.; Petro, B. J.; Vannucci, A. K.; Evans, D. H.; Glass, R. S.; Lichtenberger, D. L. J. Organomet. Chem. 2009, 694, 2681−2699. (c) Tschierlei, S.; Ott, S.; Lomoth, R. Energy Environ. Sci. 2011, 4, 2340−2352. (8) Gloaguen, F.; Lawrence, J. D.; Rauchfuss, T. B. J. Am. Chem. Soc. 2001, 123, 9476−9477. (9) Gloaguen, F.; Rauchfuss, T. B. Chem. Soc. Rev. 2009, 38, 100− 108. (10) (a) Singh, P. S.; Rudbeck, H. C.; Huang, P.; Ezzaher, S.; Eriksson, L.; Stein, M.; Ott, S.; Lomoth, R. Inorg. Chem. 2009, 48, 10883−10885. (b) Mebi, C. A.; Noll, B. C.; Gao, R.; Karr, D. Z. Anorg. Allg. Chem. 2010, 636, 2550−2554. (c) Topf, C.; Monkowius, U.; Knör, G. Inorg. Chem. Commun. 2012, 21, 147−150. (11) (a) Wright, R. J.; Lim, C.; Tilley, T. D. Chem. Eur. J. 2009, 15, 8518−8525. (b) Charreteur, K.; Kdider, M.; Capon, J.-F.; Gloaguen, F.; Petillon, F. Y.; Schollhammer, P.; Talarmin, J. Inorg. Chem. 2010, 49, 2496−2501. (12) (a) Capon, J.-F.; Gloaguen, F.; Schollhammer, P.; Talarmin, J. J. Electroanal. Chem. 2004, 566, 241−247. (b) Capon, J. F.; Gloaguen, F.; Schollhammer, P.; Talarmin, J. J. Electroanal. Chem. 2006, 595, 47−52. (c) Felton, G. A. N.; Vannucci, A. K.; Chen, J.; Lockett, L. T.; Okumura, N.; Petro, B. J.; Zakai, U. I.; Evans, D. H.; Glass, R. S.; Lichtenberger, D. L. J. Am. Chem. Soc. 2007, 129, 12521−12530. (13) Schwartz, L.; Singh, P. S.; Eriksson, L.; Lomoth, R.; Ott, S. C. R. Chim. 2008, 11, 875−889. (14) (a) Streich, D.; Astuti, Y.; Orlandi, M.; Schwartz, L.; Lomoth, R.; Hammarstrom, L.; Ott, S. Chem. Eur. J. 2010, 16, 60−63. (b) Wang, F.; Wang, W.-G.; Wang, H.-Yan; Si, G.; Tung, C.-H.; Wu, L.-Z. ACS Catal. 2012, 2, 407−416. (15) Kumar, B.; Beyler, M.; Kubiak, C. P.; Ott, S. Chem. Eur. J. 2012, 18, 1295−1298. (16) King, R. B. J. Am. Chem. Soc. 1963, 85, 1584−1587. (17) Cabeza, J. A.; Martinez-Garcia, M. A.; Riera, V.; Ardura, D.; Garcia-Granda, S. Organometallics 1998, 17, 1471−1477. (18) Hasan, M. M.; Hursthouse, M. B.; Kabir, S. E.; Malik, K. M. A. Polyhedron 2001, 20, 97−101. (19) Nomura, M.; Sasao, T.; Hashimoto, T.; Sugiyama, T.; Kajitani, M. Inorg. Chim. Acta 2010, 363, 3647−3653. (20) Darensbourg, M. Y.; Lyon, E. J.; Zhao, X.; Georgakaki, I. P. Proc. Natl. Acad. Sci. U.S.A. 2003, 100, 3683−3688. (21) An example of scan rate studies upon a two-electron reduction whose peak height does not vary with scan rate: Donovan, E. S.; Felton, G. A. N. J. Organomet. Chem. 2012, 711, 25−34. (22) Felton, G. A. N.; Petro, B. J.; Glass, R. S.; Lichtenberger, D. L.; Evans, D. H. J. Am. Chem. Soc. 2009, 131, 11290−11291. (23) Wright, R. J.; Zhang, W.; Yang, X.; Fasulo, M.; Tilley, T. D. Dalton Trans. 2012, 41, 73−82. (24) Felton, G. A. N.; Glass, R. S.; Lichtenberger, D. L.; Evans, D. H. Inorg. Chem. 2006, 45, 9181−9184. (25) Costentin, C.; Drouet, S.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 2012, 134, 11235−11242. (26) (a) Baker-Hawkes, M. J.; Billig, I. E.; Gray, H. B. J. Am. Chem. Soc. 1966, 88, 4870−4875. (b) Wharton, E. J.; McCleverty, J. A. J. Chem. Soc. A 1969, 2258−2266.
substitution upon a benzenedithiolate bridge yields an electrochemically reversible species. These easier to reduce, electrochemically reversible electrocatalysts may lead to future photon-driven proton-reduction approaches. Experimental Section. Preparation of [Fe2(μ-1,2-S2C6H34-CH3(CO)6] (1) [Fe2(μ-1,2-S2C6H4)(CO)6] (2), and [Fe2(μ-1,2S2C6H2-3,6-Cl2)(CO)6] (3). The syntheses were carried out according to literature methods and matched to published characterization data for 1,18 2,17 and 3.13 Preparation of [Fe2(μ-1,2-S2C6H-3,4,6-Cl3(CO)6] (4). This complex was prepared from 3,4,6-trichloro-1,2-dithiophenol26 (0.1 g, 0.41 mmol) and triiron dodecacarbonyl (0.21 g, 0.42 mmol) in 20 mL of tetrahydrofuran at reflux for 70 min. Silica chromatography (pentane) afforded an orange solid, which was recrystallized at −20 °C from dichloromethane and diethyl ether (3/1) to yield 4 as thin red platelike crystals suitable for X-ray crystallography. Yield: 0.051 g (24%). Anal. Calcd for C12HCl3Fe2O6S2: C, 27.53; H, 0.19; S, 12.23. Found: C, 27.70; H, 0.25; S, 12.18. 1H NMR (400 MHz, ppm, CDCl3): δ 6.83. 13 C NMR (100 MHz, ppm, CDCl3): δ 206.8, 132.4, 131.7, 129.5, 125.7. Infrared (hexanes, cm−1): νCO 2084, 2050, 2013. Preparation of [Fe2(μ-1,2-S2C6Cl4)(CO)6] (5). This complex was prepared from 3,4,5,6-tetrachloro-1,2-dithiophenol25 (0.1 g, 0.36 mmol) and triiron dodecacarbonyl (0.18 g, 0.37 mmol) in 20 mL of tetrahydrofuran at reflux for 90 min. Silica chromatography (tetrahydrofuran) afforded a red-orange solid which was recrystallized at room temperature from chloroform to yield 5 as dark red crystals suitable for X-ray crystallography. Yield: 0.069 g (35%). Anal. Calcd for C12Cl4Fe2O6S2: C, 25.83; S, 11.47. Found: C, 25.92; S, 11.55. 13C NMR (100 MHz, ppm, CDCl3): δ 206.7, 136.0, 128.5, 125.8. Infrared (hexanes, cm−1): νCO 2084, 2051, 2014.
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ASSOCIATED CONTENT
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AUTHOR INFORMATION
* Supporting Information S
Text, tables, figures, and CIF files giving further experimental details, cyclic voltammograms, IR and X-ray crystallographic details (for 4 and 5). This material is available free of charge via the Internet at http://pubs.acs.org.
Corresponding Author
*E-mail:
[email protected]. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS G.A.N.F. thanks Oakland University for start-up funds and a 2010 faculty research fellowship. E.S.D. acknowledges a Kenny graduate fellowship. J.J.M. and E.S.D. acknowledge Oakland University Provost research awards. Instrumentation grants NSF-CHE-0821487 (NMR) and NSF-CHE-1048719 (MS) are gratefully acknowledged.
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REFERENCES
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