Synthesis of Diiron Hydrogenase Mimics Bearing Hydroquinone and

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Organometallics 2010, 29, 5330–5340 DOI: 10.1021/om100396j

)

Synthesis of Diiron Hydrogenase Mimics Bearing Hydroquinone and Related Ligands. Electrochemical and Computational Studies of the Mechanism of Hydrogen Production and the Role of O-H 3 3 3 S Hydrogen Bonding† Jinzhu Chen,‡ Aaron K. Vannucci, Charles A. Mebi,§ Noriko Okumura, Susan C. Borowski, Matthew Swenson, L. Tori Lockett, Dennis H. Evans,*,^ Richard S. Glass,* and Dennis L. Lichtenberger* )

Department of Chemistry and Biochemistry, The University of Arizona, Tucson, Arizona 85721-0041. ‡ Present address: Brookhaven National Laboratory, Upton, NY. §Present address: Department of Physical Sciences, Arkansas Tech University, Russellville, AR. Present address: School of Pharmacy, Kinjo Gakuin Univeristy, Nagoya, Japan. ^Present address: Department of Chemistry, Purdue University, West Lafayette, IN. Received April 30, 2010

A new synthetic method for annulating hydroquinones to Fe2S2(CO)6 moieties is reported. Piperidine catalyzed a multistep reaction between Fe2(μ-SH)2(CO)6 and quinones to afford bridged adducts in 26-76% yields. The hydroquinone adducts undergo reversible two-electron reductions. In the presence of acetic acid, hydrogen is produced catalytically with these adducts at potentials more negative than that of the initial reversible reduction. Spectroscopic studies suggest the presence of intramolecular hydrogen bonding between the phenolic OH groups and the adjacent sulfur atoms. Computations, which are in good agreement with the electrochemical studies and spectroscopic data, indicate that the hydrogen bonding is most important in the reduced forms of the catalysts. This hydrogen bonding lowers the reduction potential for catalysis but also lowers the basicity and thereby the reactivity of the catalysts.

Introduction [FeFe]-Hydrogenase is a remarkably efficient catalyst for the production of H2. The active site of the [FeFe]-hydrogenase enzyme, shown schematically in 1, contains a 2Fe2S cluster in a butterfly arrangement.1,2 A surprising feature of the active site of the enzyme is its organometallic character with carbonyl and cyanide ligands. Long before the composition of the active site was recognized, Seyferth had characterized a diverse range of chemistry of 2Fe2S carbonyl clusters that serve as minimal models of the active site.3-6 The active site is linked to a 4Fe4S cubane cluster through a cysteinyl sulfur. The cubane cluster serves as a redox center for shuttling electrons to the active site. Also found at the active site is a bridging moiety linking the S atoms of the † Part of the Dietmar Seyferth Festschrift. In thanks to Dietmar Seyferth for his contributions to organometallic chemistry, for his guidance of this journal since its inception, and for his mentoring to me personally (D.L.L.). *To whom correspondence should be addressed. E-mail: dlichten@ email.arizona.edu (D.L.L.); [email protected] (D.H.E.); rglass@ u.arizona.edu (R.S.G.). (1) Peters, J. W.; Lanzilotta, W. N.; Lemon, B. J.; Seefeldt, L. C. Science 1998, 282, 1853–1858. (2) Nicolet, Y.; Piras, C.; Legrand, P.; Hatchikian, C.; FontecillaCamps, J. C. Struct. Fold Des. 1999, 7, 13–23. (3) Seyferth, D.; Henderson, R. S. J. Am. Chem. Soc. 1979, 101, 508–509. (4) Seyferth, D.; Henderson, R. S.; Fackler, J. P., Jr.; Mazany, A. M. J. Organomet. Chem. 1981, 213, C21–C25. (5) Seyferth, D.; Song, L.-C.; Henderson, R. S. J. Am. Chem. Soc. 1981, 103, 5103–5107.

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2Fe2S cluster. The X group of the bridging moiety is generally believed to be NH.7-10 The NH group may act as an internal pendant base or as part of a proton relay chain.11,12 In addition, the amine functionality also provides a site for hydrogen bonding with the surrounding protein. This hydrogen bonding may help stabilize the active site during catalysis.

Inspired by the active site of this enzyme, we13 and others14 have studied simple 2Fe2S complexes in an effort to find an efficient and practical catalyst for H2 production. For example, (μ-1,2-benzenedithiolato)[Fe(CO)6] (2a) was r 2010 American Chemical Society

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Organometallics, Vol. 29, No. 21, 2010

found to catalyze H2 generation from weak acids in acetonitrile on electrochemical reduction.13 The mechanism for this process and structures for the intermediates were determined by electrochemical and computational analysis. These studies provide a rational basis for designing new, more effective catalysts. Substituents on the benzene ring of 2a may be introduced to improve catalysis. In particular, hydroquinones 3 appeared promising for two reasons. The reduction of 2a, a key step in the mechanism for H2 production, results in cleavage of an Fe-S bond, resulting in the buildup of negative charge on sulfur. Adjacent phenolic groups may favor this process by hydrogen bonding to the negatively charged sulfur. Alternatively, electronwithdrawing substituents may also promote Fe-S bond cleavage. Studies with electron-withdrawing 3,6-dichloro substituents as well as aza groups, in which the benzene ring was replaced by a quinoxaline moiety, have recently been reported.15 Comparable studies with the pyridine annulated complex 14 are reported here. However, unlike the case for chloro and aza substituents, there is a second reason for studying complexes with -OH substituents attached to the benzene ring. That is, phenolic -OH groups also have an overall electron-donating electronic effect to the phenyl ring. This electron donation may increase the basicity of the corresponding dianion obtained on twoelectron reduction. Since protonation of this dianion is the rate-determining step13 in catalysis by 2a, increasing its basicity would improve catalysis. Adams and Miao16 first synthesized 3a and its mono- and disubstituted triphenylphosphine complexes and reported their electrochemical behavior. This paper reports a new and advantageous synthesis of 3a and related compounds and a reinvestigation of the electrochemistry of 3a and the ability of 3a and its analogues to catalyze H2 production from weak acids on electrochemical reduction.

Results and Discussion Synthesis. Compound 3a was synthesized previously16 in 53% yield by irradiation of 1,4-benzoquinone (4a) and (6) Cowie, M.; DeKock, R. L.; Wagenmaker, T. R.; Seyferth, D.; Henderson, R. S.; Gallagher, M. K. Organometallics 1989, 8, 119–132. (7) Nicolet, Y.; de Lacey, A. L.; Vernede, X.; Fernandez, V. M.; Hatchikian, C.; Fontecilla-Camps, J. C. J. Am. Chem. Soc. 2001, 123, 1596–1601. (8) Nicolet, Y.; Cavazza, C.; Fontecilla-Camps, J. C. J. Inorg. Biochem. 2002, 91, 1–8. (9) Silakov, A.; Wenk, B.; Reijerse, E.; Lubitz, W. Phys. Chem. Chem. Phys. 2009, 11, 6592–6599. (10) Pandey, A. S.; Harris, T. V.; Giles, L. J.; Peters, J. W.; Szilagyi, R. K. J. Am. Chem. Soc. 2008, 130, 4533–4540. (11) Li, P.; Wang, M.; Chen, L.; Lui, J.; Zhao, Z.; Sun, L. Dalton Trans. 2008, 1919–1926. (12) Barton, B. E.; Olsen, M. T.; Rauchfuss, T. B. J. Am. Chem. Soc. 2008, 130, 16834–16835. (13) 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. (14) Gloaguen, F.; Rauchfuss, T. R. Chem. Soc. Rev. 2009, 348, 100– 108. Tard, C.; Pickett, C. J. Chem. Rev. 2009, 109, 2245–2274. 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. Capon, J.-F.; Gloaguen, F.; Petillon, F. Y.; Schollhammer, P.; Talarmin, J. Coord. Chem. Rev. 2009, 253, 1476–1494. (15) Schwartz, L.; Singh, P. S.; Eriksson, L.; Lomoth, R.; Ott, S. C. R. Chim. 2008, 11, 875–889. (16) Adams, R. D.; Miao, S. Inorg. Chem. 2004, 43, 8414–8426. (17) Seyferth, D.; Henderson, R. S. J. Organomet. Chem. 1981, 218, C34–C36.

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Scheme 1. Addition of 2-Cyclohexenone to 6, Resulting in a,e and e,e Isomers

Table 1. Reaction of Dithiol 6 with 1,4-Quinones in the Presence of Piperidine quinone

product

yield (%)a

4a 4b 4c 4d 4e 4f 8 10

3a 3b 3c 3d 3e 3f 9 11

64 70 45 26 50 53 76 57

a

Yield of isolated pure product.

disulfide 5. Another expeditious route, however, to compound 3a was discovered by serendipity in the present work. 2-Cyclohexenone, like methyl acrylate, methyl vinyl ketone, and acrylonitrile, undergoes a 1,4-conjugate addition to dithiol 6 in the presence of piperidine (Scheme 1) to afford 2:1 adducts as mixtures of a,e and e,e isomers.17-19 Since it is well-known that thiols and thiolates can add conjugatively to the carbon-carbon double bond of quinones,20-25 it was surmised that the reaction of dithiol 617 or dithiolate 726,27 with 1,4-benzoquinone would proceed analogously to that of 2-cyclohexenone to provide 2:1 adducts as mixtures of a,e and e,e isomers. Reaction of the dianion 7 with 1,4-benzoquinone did not result in adduct formation but rather oxidation of the dianion 7 to disulfide 5 by the quinone. Surprisingly, reaction of dithiol 6 with 1,4-benzoquinone in the presence of piperidine did not afford a mixture of stereoisomeric 2:1 adducts but rather 3a in 64% isolated yield. Similarly, reaction of the 2-substituted 1,4-benzoquinones 4b-f with dithiol 6 in the presence of piperidine proceeded in an analogous manner to provide the corresponding hydroquinone adducts 3b-f in the yields reported in Table 1. 1,4-Naphthoquinone (8) reacted similarly to give adduct 9 in 76% isolated yield. Treatment of 1,4-anthraquinone (10) analogously led not to the corresponding hydroquinone adduct but rather quinone adduct 11 in 57% (18) Seyferth, D.; Womack, G. B. J. Am. Chem. Soc. 1982, 104, 6839– 6841. (19) Seyferth, D.; Womack, G. B.; Henderson, R. S. Organometallics 1986, 5, 1568–1575. (20) Finley, K. T. In The Chemistry of the Quinonoid Compounds; Patai, S., Rappoport, Z., Eds.; Wiley: New York, 1988; Vol. 2, Part 1, pp 533-717. (21) Kutyrev, A. A. Tetrahedron 1991, 47, 8043–8065. (22) Pacsial-Ong, E. J.; McCarley, R. L.; Wang, W.; Strongin, T. M. Anal. Chem. 2006, 78, 7577–7581. (23) Yonge, L.; Gracheva, S.; Wilkins, S. J.; Livingstone, C.; Davis, J. J. Am. Chem. Soc. 2004, 126, 7732–7733. (24) Watanabe, N.; Dickinson, D. A.; Liu, R.; Forman, H. J. Methods Enzymol. 2004, 378, 319–340. (25) Bratton, S. B.; Lau, S. S.; Monks, T. J. Chem. Res. Toxicol. 1997, 10, 859–865. (26) Seyferth, D.; Henderson, R. S.; Song, L. C. J. Organomet. Chem. 1980, 192, C1–C5. (27) Seyferth, D.; Henderson, R. S.; Song, L. C. Organometallics 1982, 1, 125–133.

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isolated yield. Compound 12 was produced from 11 by electrolysis.

Figure 1. Comparison of simulations (circles) with background-corrected voltammograms (full curves) for 1.00 mM 2a, 1.03 mM 3a, and 1.00 mM 3f. Conditions: acetonitrile with 0.10 M NBu4PF6, glassy-carbon electrode, scan rate 0.10 V/s, room temperature.

Note that this new method provided 3a in a somewhat higher yield than the photochemical method (64 vs 53%) and the preparation was less prone to decomposition in air by this new method. To further compare the two synthetic methods, 9 was prepared photochemically in 43% yield and 11 was prepared photochemically in 19% yield. Comparison of these photochemical yields with those obtained by the new method (76 and 57%, respectively) reveals that the new method provides apparently higher yields, in our hands, than the photochemical synthesis, but neither method was fully optimized. The photochemical reaction, however, requires fewer steps than the new method. Conjugate addition of the SH moieties of dithiol 6, in the presence of piperidine, to R,β-unsaturated ketones, nitriles, and esters results in mixtures of stereoisomers a,e and e,e of 2:1 adducts, as illustrated for 2-cyclohexenone in Scheme 1. Formation of 3, 9, and 11, however, requires that the initial conjugate addition of the thiolates of 6 to the corresponding quinone provide the axial (a) rather than the equatorial (e) isomer. Formation of 1:1 adducts with initial formation of the axial isomer has been reported before15,19 in the reaction of β,β-disubstituted R,β-unsaturated ketones, resulting in the formation of bridged adduct 13. This anomalous result was ascribed to steric hindrance in the monoadduct by the tertiary carbon center on the nearby sulfur center, preventing addition of a second molecule. Formation of 1:1 bridged products was also found15,19 for the reaction of dithiol 6 with activated acetylenes. This was ascribed to two consecutive conjugate additions. Following the initial conjugate addition of dithiol 6 to the quinone, there must be an oxidation of the organic ligand to a sulfur-substituted quinone to enable (28) The corresponding quinoxaline derivative has been reported.15

Scheme 2. Proposed Pathway for Conversion of 6 and Quinone to the a,a-Bridged Product

the second conjugate addition to afford the observed a,abridged product. This is outlined in Scheme 2. The starting quinone 4 serves as the oxidizing agent.

To separate hydrogen-bonding effects from substituent effects in the electrochemistry and catalytic H2 production in hydroquinone 4a, the dimethoxy derivative 2b was desired. Hydroquinone 3a was O-methylated to give 2b by treatment with potassium carbonate and methyl iodide.

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Table 2. Simulation Parameter Values for Reduction of 3a-f, 9, 12, 2a,b, and 14a compd

E1

E2

Eov

D/cm2 s-1

3a 3b 3c 3d 3e 3f 9b 12b 2b 2a 14

-1.400 -1.460 -1.407 -1.403 -1.326 -1.298 -1.522

-1.158 -1.170 -1.159 -1.146 -1.119 -1.140 -1.160

1.45 10-5 1.35 10-5 1.27 10-5 1.25 10-5 1.25 10-5 1.25 10-5 8.43 10-6

-1.389 -1.444 -1.238

-1.296 -1.212 -1.236

-1.279 -1.315 -1.283 -1.274 -1.222 -1.219 -1.341 -1.38 -1.342 -1.328 -1.237

c

c

c

1.40 10-5 1.45 10-5 1.45 10-5

Figure 2. Calculated structures of the dianions 2a2-, 2b2-, and 3a2-, illustrating that all three compounds undergo large geometric changes upon 2e reduction.

a All potentials are reported as V vs the Fcþ/Fc potential in the solvent that was used (acetonitrile unless otherwise indicated). Conditions: room temperature, scan rate 0.1 V/s. Diffusion coefficients were considered to be identical for the neutral compound, anion, and dianion. Eov is the standard potential for the overall two-electron reduction defined by (Eo1 þ Eo2)/2. Standard electron-transfer rate constants were 0.10 cm/s for the first step of reduction and 0.028 cm/s or the second (except for 14, where 0.001 and 0.13 cm/s were used). Electron-transfer coefficients were 0.50 for both steps. The extent of potential inversion, E1 - E2, is difficult to evaluate accurately (compare results for 2a with those in ref 13). The overall potential, Eov, is much more reproducible. b In dichloromethane. The solution was prepared by electrolysis of (μ-1,4-naphthoquinone-2,3-dithiolato)Fe2S2(CO)6 (for 9) or 11 (for 12) in the presence of 2 equiv of acetic acid. c The purity of 12 formed by electrolysis was low, making the determination of individual standard potentials and the diffusion coefficient of questionable accuracy. The potentials were inverted.

Electrochemistry. The effectiveness of the hydroquinonecontaining catalysts 3a-f, 9, and 12 toward catalyzing the reduction of protons from acetic acid was assessed. Also studied were the dimethoxy derivative 2b, the pyridine-containing complex 14,29 and a previously studied complex with a benzenedithiolate ligand, 2a.13 In the absence of added acid, the hydroquinone-containing catalysts undergo reversible two-electron reductions. Examples of the cyclic voltammograms of 3a,f, along with 2a, are given in Figure 1, where simulations are compared to experimental results. The reductions occur with potential inversion: i.e., introduction of the second electron occurs more easily than the first; this may be seen by examining the parameter values used in the simulations (Table 2). For compounds 3a-f, the effect of ring substitutions on the reduction potentials is modest but outside experimental error. Earlier computational work13 indicated that significant structural changes accompany the reduction of 2a. These changes include a lengthening of an Fe-S bond with its eventual full cleavage in the dianion as well as movement of a terminal CO to a bridging position. It is presumed that analogous structural changes occur upon reduction of 3a-f, 2b, and 14, since these complexes exhibit a striking similarity in electrochemical behavior with that of 2a. This expectation was supported by calculations for 2b and 3a, where the computed structures for the dianions are quite analogous to that seen for 2a (Figure 2). Specifically, breakage of an ironsulfur bond and movement of a terminal CO to a bridging position occurs. In the case of 2a, these structural changes activate the system for participation in a catalytic cycle leading to the reduction of weak acids to form H2. The present compounds (29) Petro, B. J.; Vannucci, A. K.; Lockett, L. T.; Mebi, C.; Kottani, R.; Gruhn, N. E.; Nichol, G. S.; Goodyer, P. A. J.; Evans, D. H.; Glass, R. S.; Lichtenberger, D. L. J. Mol. Struct. 2008, 890, 281–288.

Figure 3. Voltammograms of about 1.0 mM catalyst plus increasing concentrations of acetic acid (0, 1, 2, 3, 5, 7, 10, 15, 20, 25, 35, and 50 mM). Other conditions are as in Figure 1. A compensation of 140 Ω of solution resistance was applied electronically.

behave quite similarly. Figure 3 contains voltammograms for 3f and 2b in the presence of increasing quantities of acetic acid (pKa = 22.330 in acetonitrile). As can be seen, in the presence of excess acetic acid, a catalytic peak appears at potentials more negative than the initial reduction potentials of the compounds, and the height of the catalytic peak increases with increasing concentrations of the acid. Also noticeable from Figure 3, the peak current for the catalytic peak of 2b is much greater than the catalytic peak current for 3f, and the onset of catalytic peak current for 3f is at a much less negative potential than for 2b even though the initial two-electron reduction is at only slightly less negative potential. For a more complete comparison, voltammograms of 2a,b and 3a,c-f, obtained in the presence of 50 mM acetic (30) Izutsu, K. In Acid-Base Dissociation Constants in Dipolar Aprotic Solvents; Blackwell Scientific Publishers: Oxford, U.K., 1990.

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Figure 4. Voltammograms of about 1.0 mM catalysts with 50 mM acetic acid. Other conditions are as in Figure 3.

Chen et al.

Figure 6. Examples of fits of background-corrected voltammograms (full curves) to simulations (symbols) for 1.05 mM 3a in the presence of the indicated concentrations of acetic acid at a scan rate of 0.30 V/s. See Table 3 for values of the simulation parameters.

acid, are displayed in Figure 4. The results in Figures 3 and 4 indicate that there are large differences in the heights of the catalytic peaks and significant differences in the positions of these peaks on the potential axis for the hydroquinonecontaining catalysts (3a-f) compared to 2a,b. There are two measures of the efficiency of this catalysis: the magnitude of the catalytic current and the potential required to produce the current. For the former, we take the peak current for a catalyst in the presence of a given acetic acid concentration (50 mM), Ip,cat, and for the latter we choose the potential where the catalytic current reaches half its peak value, Ep/2,cat. A plot of these quantities for 2a,b, 3a, c-f, and 14 is given in Figure 5. Note the standard potential for reduction of acetic acid, -1.46 V, is indicated on the graph.31 It is immediately apparent that those catalysts for which the catalytic peak appears at the more negative potentials are the catalysts with the largest peak currents. Defining the overpotential of catalysis as Ep/2,cat - E leads to overpotentials from -0.54 V for 2b to -0.37 for 3f. However, the catalytic peak current for the lowest-overpotential catalyst, 3f, is only about 25% of that seen with the high-overpotential catalyst 2b. Thus, low overpotential is

achieved at the cost of low rates of catalysis. The arbitrary extrapolation of the plot in Figure 5 suggests that, in this family of catalysts, there would be almost no catalyzed reduction of acetic acid near zero overpotential. A related observation was made in our studies of 2a9 and [(η5C5H5)Fe(CO)2]2,32 where it was shown that reductions of acids with different pKa values and the same catalyst led to lower overpotentials being associated with smaller catalytic currents. Since the catalysts in this study exhibit reduction chemistry similar to that of 2a, the mechanism of the catalyzed reduction of protons from acetic acid for these catalysts should be analogous to the proposed catalytic mechanism for 2a. To test this point, we have simulated results for one of the catalysts, 3a, and acetic acid and find excellent agreement (see Figure 6). The reaction scheme is given by reactions 1-7 in Table 3, in which Cat refers to 3a and HA is acetic acid. Table 3 also includes the values of the standard potentials of reactions 1 and 2 computed by our density functional theory model (vide infra). The agreement of the computed overall potential, Eov, with experiment (-1.32 V calculated vs -1.30 V experimental) is very good and is the most meaningful, as this is the quantity that can be obtained reliably from the voltammograms (see footnote a of Table 2). Calculations also indicate that the structures of CatH- and CatH2- for 3a are analogous to those reported for 2a.13 Despite the similar calculated structures and catalytic reduction mechanisms of the benzenedithiolato complex 2a and the hydroquinonedithiolato complex 3a, 2a is a much better catalyst, producing more than twice the catalytic current for hydrogen production. The dimethoxybenzenedithiolato complex 2b is also a much better catalyst than 3a, despite the similar inductive and hyperconjugative effects of the -OCH3 and -OH substituents. Internal hydrogen bonding of the phenolic -OH groups to the S atoms in molecules of type 3 may be responsible for the differences in efficiency of these catalysts. Evidence for internal hydrogen bonding and its effects on catalysis is discussed in the following sections.

(31) Felton, G. A. N.; Glass, R. S.; Lichtenberger, D. L.; Evans, D. H. Inorg. Chem. 2006, 45, 9181–9184. Felton, G. A. N.; Glass, R. S.; Lichtenberger, D. L. Inorg. Chem. 2007, 46, 5126.

(32) Felton, G. A. N.; Vannucci, A. K.; Okumura, N.; Lockett, L. T.; Evans, D. H.; Glass, R. S.; Lichtenberger, D. L. Organometallics 2008, 27, 4671–4679.

Figure 5. Catalytic peak current as a function of the half-peak potential of the catalytic peak. Conditions: 1.0 mM catalyst plus 50 mM acetic acid in acetonitrile. Other conditions are as in Figure 3.

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Table 3. Reactions and Values of Parameters Used in the Simulation of the Catalytic Reduction of Acetic Acid using 3aa no.

reacn

params

1

Cat þ e - h Cat• Cat• - þ e- h Cat2 CatH - þ e- h CatH2 - • 2Cat• - h Cat þ Cat2 Cat2 - þ HA h CatH - þ A CatH•2 - þ HA h Cat2 - þ A - þ H2 HA þ A - h HA2 -

E1, ks,1, R1

2 3 4 5 6 7 no. 1 2 4

E2, ks,2, R2 E3, kf,3, kb,3 E4, kf,4, kb,4 E5, ks,5, R5 E6, kf,6, kb,6 E7, kf,7, kb,7 R

reacn

E/V

ks/cm s-1

Cat þ e - h Cat• Cat• - þ e - h Cat2 CatH - þ e - h CatH•2 -

-1.39 (-1.61)

0.2

0.5

-1.20 (-1.02)

0.1

0.5

-1.92 (-2.02)

0.05

0.5

no.

reacn

K

3

2Cat• - h Cat þ Cat2 Cat2 - þ HA h CatH - þ A CatH•2 - þ HA h Cat2 - þ A - þ H2 HA þ A - h HA2 -

5.0 10-4

103

0.05

2.0 10

4.0 105

0.10

b

b

1000

1.0 10

4 6 7

kf

kb 2.0 106 4

8

1.0 105

a For simulations like those shown in Figure 6. Values in parentheses are computed values of standard potentials. The same parameters were used to fit data for 0.1, 0.3, 0.5, and 1.0 V/s and for 5, 10, 20, and 50 mM acetic acid. Ei, ks,i, and Ri are the standard potentials (V vs ferrocene), standard rate constants (cm/s), and electron transfer coefficients of electrode reactions i. Kj, kf,j, and kb,j are the equilibrium constant, forward rate constant, and reverse rate constant of chemical reaction j. Units of rate constants of chemical reactions: s-1 for first-order reactions and M-1 s-1 for second-order reactions. The value found for the equilibrium constant of the homoconjugation reaction (7) (1000 M-1) may be compared to the reported30 value, 6000 M-1. The total solution resistance to be compensated was found to be 180 Ω. Of this, 120 Ω was compensated by the potentiostat and 60 Ω was included in the simulations.Electrode area: 0.08 cm2. Diffusion coefficients of all “Cat” species: 1.5 10-5 cm2/s; for HA, A- and HA2-: 3.0 10-5 cm2/s; for H2: 5.0 10-5 cm2/s. b By computation we could find no evidence for an intermediate in reaction 6 that contained bound dihydrogen, CatHH•-. However, due to software limitations, it was necessary to treat the overall reaction CatH•2- þ HA h Cat2- þ A- þ H2 as two fast sequential reactions, CatH•2- þ HA h CatHH•- and CatHH•- h Cat•- þ H2.

Spectroscopy. To evaluate whether internal hydrogen bonding occurs in the hydroquinone-containing compounds 3, the photoelectron and IR spectra of 3a were compared to those of the dimethoxy derivative 2b. Valence photoelectron spectra were collected for both compounds and are shown in Figure 7. The first broad ionization band starting near 7.5 eV ionization energy and extending to roughly 9 eV contains mainly the iron-based d-electron ionizations. Ionizations starting just beyond 9 eV and continuing to higher energy are associated with ionizations of the dithiolate ligands and show the greatest variation between the two compounds. The substantial sulfur character that gives rise to these ionizations is shown by the considerable decrease in ionization intensity relative to the iron d-based region when the He II spectrum is compared to the He I spectrum.29,33 The experimental adiabatic ionizations, observed as the low-energy onset of ionization intensity, were compared to calculated values. The arrows on the low-energy side of the spectra denote the calculated adiabatic ionization energies, which closely match the experimental values. This agreement indicates that the electronic structure calculations are accounting for the change in energies resulting from the changes in electron configurations and geometries of these systems reasonably well, as was also found for the calculated reduction potentials discussed above. Various studies have shown that DFT methods can accurately calculate noncovalent interactions that are electrostatic in nature, such (33) Glass, R. S.; Gruhn, N. E.; Lorance, E.; Singh, M. S.; Stessman, N. Y. T.; Zakai, U. I. Inorg. Chem. 2005, 44, 5728–5737. (34) Zhao, Y.; Truhlar, D. G. J. Chem. Theory Comput. 2006, 2, 1009–1018.

Figure 7. He I (solid line) and He II (dashed line) photoelectron spectra of 3a and 2b. Arrows indicate adiabatic energies calculated by the DFT model. The He II intensity is scaled to match the He I intensity along the leading ionization edge.

as hydrogen bonding.34-36 The strong agreement between the calculated and experimental redox potentials, ionization energies, IR stretching frequencies, and other aspects of this study gives credence to the ability of these calcu(35) Riley, K. E.; Hobza, P. J. Phys. Chem. A 2007, 111, 8257–8263.

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Figure 9. Three possible low-energy conformers of 3a and 2b (X = H, Me, respectively).

Figure 8. Comparison of the experimental IR spectra in the carbonyl stretching region (top) with the combined theoretical simulations for all three conformers (bottom) for compounds 3a and 2b. The positions and relative intensities (height of the colored lines) of the theoretical stretching frequencies are schematically shown as the colored lines on the bottom spectra (see the Supporting Information).

lations to properly model this intramolecular hydrogen bonding. Evidence of hydrogen bonding in 3a is seen in both the adiabatic and the sulfur-based ionizations. Hydrogen bonding from the hydroquinone ligand to the sulfurs results in stabilizing the complex as a whole. The adiabatic ionization energy of complex 3a is roughly 0.14 eV higher than that of 2b. Further, because hydrogen bonding is between the phenolic OHs and the sulfur atoms, it is expected that the sulfur-based ionizations should be stabilized in 3a relative to 2b. As can be seen in Figure 7, the sulfur-based ionizations, located primarily in the energy range of 9-11 eV, are stabilized by nearly 0.4 eV in 3a compared to 2b. This dramatic shift of ionization energy has not been seen in similar compounds that do not have the possibility of hydrogen bonding,29,33 giving further indication that internal hydrogen bonding between the phenolic hydrogen atoms and the sulfur atoms has a large effect on the 2Fe2S core of these compounds. Computations also indicate that the sulfur lone pair orbitals for 3a, in a conformation that would allow hydrogen bonding, are stabilized by roughly 0.3 eV compared to sulfur lone pair orbitals of 3a in a conformation that lacks any intramolecular hydrogen bonding. This stabilization of the lone pair orbitals should lead to broadening of the sulfur peak in the photoelectron spectrum, which is indeed observed in the spectrum of 3a shown in Figure 7. Conversely, the calculated energies of the sulfur lone pair orbitals for various conformations of 2b differ by less than 0.1 eV, which matches y, J.; Hobza, P. J. Chem. Theory (36) Riley, K. E.; Pito nak, M.; Cern Comput. 2010, 6, 66–80.

the lesser extent of broadening in the sulfur lone pair peaks in the experimental photoelectron spectrum of 2b. For evidence that the hydrogen bonding also occurs for the molecules in solution, IR spectra in heptane solution were obtained for both compounds, and the spectra in the range of the carbonyl stretching frequencies are compared to calculated analytical frequencies in Figure 8. The carbonyl stretching frequencies give insight into the amount of stabilization of the 2Fe2S core, because the carbonyl stretching frequencies are dependent on the electron richness of the metal center and the extent of back-bonding to the carbonyls. For both 2b and 3a, three different conformations were calculated to obtain the calculated IR spectra. The conformations, illustrated in Figure 9, pertain to differences in the possible orientations of the phenolic OH and OMe groups. Furthermore, these conformers should equilibrate readily because the barrier for their interconversion should be low and comparable to the barriers of rotation for 1,4-dihydroxybenzene and 1,4-dimethoxybenzene.37 With respect to compound 3a, the down-up and down-down conformers can participate in hydrogen-bonding interactions between the hydrogen atoms from the phenolic OH groups and the nearby sulfur atoms. The hydrogen-bonding interactions in 3a should stabilize the 2Fe2S core, which is predicted by calculations to be on the order of 3 kcal/mol. This stabilization should cause a shift in the carbonyl stretching frequencies to higher frequencies, and the shift is observed in both the experimental and calculated spectrum for 3a (Figure 8). The various conformations of 3a result in small shifts in the CO stretching frequencies, causing the observed broadening, which is reproduced very well in the calculations. For 2b, the OMe groups cannot participate in hydrogen-bonding interactions with the S atoms, and the IR spectrum of 2b lacks the broadening observed for 3a. Also, IR spectra of 2b in a large excess of acetonitrile (Supporting Information) showed no shifting or broadening in the CO peaks. The three different conformers of 2b differ by less than 1 kcal/mol. Hence, the calculated CO stretching frequencies of 2b overlap much more than the frequencies for 3a, leading to narrow peaks in the calculated spectrum, which match the narrow peaks of the experimental spectrum of 2b very well. Role of Hydrogen Bonding in Electrochemistry and Catalysis. Electrochemical results also showed the influence of intramolecular hydrogen bonding in 3a. Both the initial reduction potential and the potential for electrocatalytic production of hydrogen for 3a were shifted to less negative potentials compared to 2b. The reduction potentials for iron-sulfur clusters are known to be less negative when (37) Radom, L.; Hehre, W. J.; Pople, J. A.; Carlson, G. L.; Fateley, W. I. Chem. Commun. 1972, 308–309. Kim, K.; Jordan, K. D. Chem. Phys. Lett. 1995, 241, 39–44. (38) Ueyama, N.; Yamada, Y.; Okamura, T.; Kimura, S.; Nakamura, A. Inorg. Chem. 1996, 35, 6473–6484.

Article

Figure 10. Three different conformations for 3a2-. The hydrogen bonding between the H atom of the phenolic OH group to the S atom greatly stabilizes the compound.

hydrogen bonding between the cluster and an amide N-H35 or a phenol O-H39 is possible. The change in reduction potential is believed to be of biological significance for 2Fe2S clusters40-43 as well as 4Fe4S clusters.40,44 Calculations give insight into the stabilizing effects observed in the electrochemical experiments. The proposed hydrogen bonding in 3a is expected to have an even larger influence on the dianion compared to the neutral species because of the greater negative charge at the sulfur atoms. Hydrogen bonding in 3a2- in acetonitrile as solvent is supported by computational studies which show that the down-down conformer is the most stable conformation (Figure 10). The similar energies in the down-up and down-down conformers of 3a2can be explained through an examination of the calculated Voronoi deformation densities on the atoms. The Voronoi deformation density, which is less sensitive to the variation of computational parameters than Mulliken population analyses,45 is a measure of the electron density distribution in the compounds. For the neutral conformers of 3a, the partial charge on the S atoms ranged from -0.01 for the downdown species to -0.06 for the up-up species. Once 3a2- is formed, the S atoms no longer possess similar charges. The S atom that has cleaved from one of the Fe atoms (front S atoms in Figure 10) possesses a much more negative charge than the S atom still involved in three bonds. The more negative sulfur atom in the up-up conformation of 3a2- has a -0.24 partial charge, while the other S atom still only has a -0.02 charge. In the down-down and down-up conformations, the negative charge on the more negative S atom is reduced to -0.20 in both conformations. This reduction of charge density from the hydrogen-bonding interactions has a stabilizing effect and favors the hydrogen-bonded conformations. Hydrogen bonding in 3a2- provides an explanation of the differences in the reduction potentials of 3a compared to 2b. Substituent effects in compounds 3b-f also influence the reduction potentials. The basis for these substituent effects likely results from the interplay of electronic effects on the sulfur atoms as well as the effects on the phenolic -OHs. (39) Yang, X.; Niu, S.; Ichiye, T.; Wang, L. J. Am. Chem. Soc. 2004, 126, 15790–15794. (40) Stephens, P. J.; Jollie, D. R.; Warshel, A. Chem. Rev. 1996, 96, 2491–2513. (41) Vidakovic, M.; Fraczkiewicz, G.; Dave, B. C.; Czernuszewicz, R. S.; Germanas, J. P. Biochemistry 1995, 34, 13906–13913. (42) Denke, E.; Merbitz-Zahradnik, T.; Hatzfeld, O. M.; Snyder, C. H.; Lick, T. A.; Trumpower, B. L. J. Biol. Chem. 1998, 273, 9085– 9093. (43) Sevrioukova, I. F. J. Mol. Biol. 2005, 347, 607–621. (44) Babini, E.; Borsari, M.; Capozzi, F.; Eltis, L. D.; Luchinat, C. J. Biol. Inorg. Chem. 1999, 4, 692–700. (45) Guerra, C. F.; Handgraaf, J. W.; Baerends, E. J.; Bickelhaupt, F. M. J. Comput. Chem. 2004, 25, 189–210.


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While the relative importance of these effects is difficult to assess, the ensuing discussion focuses on the substituent effects on hydrogen bonding because of the apparent dominance of hydrogen-bonding effects by the phenolic -OHs for the difference in reduction potentials of 3a and 2b discussed above. The substituent effects, however, are complicated by the fact that the substituent is ortho to one -OH and meta to the other and ortho-substituent effects include steric as well as electronic factors. Ortho-substituent effects on the acidity of phenols have been studied experimentally and computationally,46 and acidity has been shown to correlate with hydrogen-bonding ability.47 Thus, the greater the acidity of the phenolic -OHs in 3b-f, the greater its proclivity to hydrogen bonding intramolecularly, resulting in a less negative reduction potential. Compounds 3a,c,d show the same overall reduction potentials within experimental error, as expected on the basis of comparable σm values48 and comparable pKas for phenol and o-cresol in acetonitrile.30 On the other hand, compounds 3e,f show comparable overall reduction potentials that are 57 and 60 mV, respectively, less negative than that of 3a. These less negative reduction potentials are in accord with the comparable positive σm values for Cl and Br48 and greater acidity of both o-bromophenol and m-chlorophenol compared to phenol in acetonitrile.30 Compound 3b presents a dilemma because σm48 is positive for MeO but smaller than that for Cl or Br, whereas an o-MeO is calculated to lower the acidity of phenol in the gas phase.46 It is also known47 that the OH and OMe are intramolecularly hydrogen bonded in 2-methoxyphenol. Consequently, S and O compete for the hydrogen bonding to the adjacent OH in 3b2-. It is not unreasonable that these competing factors result in the observed more negative (36 mV) overall reduction potential of 3b compared with 3a, but this complication in accounting for the results with 3b points out a limitation in the analysis. Compound 14 demonstrates the effect of moieties in which hydrogen bonding is not a factor. It is expected that the electron-withdrawing aza nitrogen would stabilize the corresponding dianion, and indeed its reduction potential is less negative than that of 2a by 91 mV. While internal hydrogen bonding was able to lower the overpotentials associated with the reduction of protons from acetic acid to catalytically produce molecular hydrogen, internal hydrogen bonding also resulted in reduced catalytic efficiency (magnitude of the catalytic peak current). The highest catalytic efficiency was associated with catalysts incapable of intramolecular hydrogen-bonding which also have the most negative half-peak potential of the catalytic process. These observations suggest that a leading cause of the changes in catalytic efficiencies among this family of catalysts (arene moieties annulated to a Fe2S2(CO)6 core) are changes in the equilibrium constant of the rate-determining step (rds), Cat2- þ CH3CO2H h CatH- þ CH3CO2-. In fact, for the two catalysts where this equilibrium constant has been evaluated through electrochemical simulations, 3a and 2b, the values of the equilibrium constants correlate with catalytic activity. The equilibrium constant corresponding to the rate-determining step for 3a was found to be 0.05, while (46) Himo, F.; Noodleman, L.; Blomberg, M. R. A.; Siegbahn, R. E. M. J. Phys. Chem. A 2002, 106, 8757–8761. (47) Korth, H.; de Heer, M. I.; Mulder, P. J. Chem. Phys. A 2002, 106, 8779–8789. (48) March, J. In Advanced Organic Chemistry, 4th ed.; Wiley: New York, 1992; p 280.

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the value for 2b, which has a much greater catalytic efficiency, was determined to be 2.0. Other factors, including the rate constants of this reaction, are also important, but the effects of the equilibrium constant may dominate the behavior. Wayner and Parker49 pointed out that there was a relationship among the pKa of a metal hydride, the homolytic M-H bond dissociation free energy of that metal hydride, the standard potential for formation of the metal anion, and the standard potential of the Hþ/H• couple. This relationship is relevant to this work because the pKa of CatH- will govern the equilibrium constant of the rds for each of the catalysts studied since the other partner in the rds is acetic acid in every case. Reactions 8-910 may be combined to obtain the desired acid dissociation constant, K2,H2Cat, as given in eq 11, where E2 has been defined earlier (Table 2), ΔGBD is the homolytic bond dissociation free energy of the metal-hydride bond in CatH-, and EH is the standard potential of the Hþ/ H• couple, which is a quantity that is independent of the nature of the catalyst. The value of ΔGBD can also be considered a constant, because it was found that for a given metal ΔGBD was constant to within 0.1-2.7 kcal/mol with a variety of different ligands in the hydrides.49 The relatively minor changes in the ligand structure in our complexes may suggest that ΔGBD varies little among them. Therefore, by taking the partial derivative of log K2,H2Cat with respect to E2, with ΔGBD and EH held constant, a simple prediction (eq 14) is obtained: the acid dissociation constant of CatHwill increase by a factor of 10 for every 59 mV change of E2 in the positive direction. Otherwise stated, reductions occurring at more negative potentials will have products that are more basic. Hence, eq 14 provides a chemical rationalization for the trend readily apparent in Figure 5, that greater catalytic efficiency requires larger overpotentials for this family of catalysts.

stant for protonation and the lower the catalytic efficiency. The arguments presented above for the role of hydrogenbonding and substituent effects on the overall reduction potentials of 3a-f and 2a,b apply here. That is, the dianions of 3a-f are greatly stabilized by hydrogen bonding between the phenolic OH groups and the S atoms of the 2Fe2S core of the catalysts. This hydrogen bonding resulted in reduced catalytic efficiencies for 3a-f compared to 2a,b. Additionally, the increased strength of the hydrogen bonding in 3e,f due to the electron-withdrawing Cl and Br substituents stabilized the 2Fe2S core, further disfavoring protonation relative to 3a,c,d. This resulted in even less catalytic efficiency for 3e,f compared with 3a,c,d. Ideally, control of hydrogen bonding in the dianion could result in improved catalytic efficiency. The hydrogen bonding already resulted in less negative reduction potentials; however, if hydrogen bonding could be removed after reduction, the increased basicity of the dianion would result in increased magnitude of the catalytic peak current.

Conclusions In an attempt to design catalysts that improve catalytic efficiency and lower catalytic overpotentials compared to 2a, catalysts of the general structure 3 were synthesized and investigated. Phenolic -OH groups adjacent to the sulfur centers in Fe2S2 complexes 3 modestly render their reduction potentials less negative and hence lowered the catalytic overpotential compared with the unsubstituted complex 2a due to internal hydrogen bonding. Electron-withdrawing aza groups also lowered the reduction potentials of 14 compared to 2a, due to an electronic effect previously reported.15 The hydroquinone complexes 3, however, are less effective catalysts for H2 production from weak acids than 2a because of the internal hydrogen bonding. The internal hydrogen bonding in 3 resulted in lowered basicity of the Cat2- species of the rate-determining step. However, the dimethoxy derivative 2b, in which hydrogen bonding is not possible, is a better catalyst than 2a owing to the increased basicity of the corresponding dianion. These results further indicate that the pKa value of the CatH- species is the most important factor for catalytic activity for this family of catalysts.

Experimental Section ln K2, H2 Cat ¼ ðF=RTÞE2 - ΔGBD =RT - ðF=RTÞEH

ð12Þ

log K2, H2 Cat ¼ ðF=ðln 10ÞRTÞE2 - ΔGBD =ðln 10ÞRT - ðF=ðln 10ÞRTÞEH

ð13Þ

D log K2, H2 Cat 1 V -1 ¼ 0:05916 DE2 at 298 K if ΔGBD is constant

ð14Þ

The underlying basis for the change in the equilibrium constant of the rate-determining step, Cat2- þ CH3CO2H h CatH- þ CH3CO2-, is the basicity of Cat2-. The more stabilized the dianion, the less favorable the equilibrium con(49) Wayner, D. D. M.; Parker, V. D. Acc. Chem. Res. 1993, 26, 287– 294.

General Considerations. All reactions were carried out under an atmosphere of prepurified nitrogen by using standard Schlenk and vacuum-line techniques. Tetrahydrofuran (THF) and hexane were purified by distillation under nitrogen from sodium/ benzophenone ketyl. Dichloromethane was distilled from calcium hydride under nitrogen. Acetonitrile was distilled once from P2O5 and then freshly distilled from CaH2 under N2 before use. Ethanol was distilled from magnesium under nitrogen. LiBEt3H (1 M in THF), trifluoroacetic acid, 1,4-benzoquinone, substituted 1,4-benzoquinones, 1,4-naphthoquinone, 1,4-anthraquinone, piperidine, and 2,3-dichloro-5,6-dicyano-1,4-benzoquinone were purchased from Aldrich Chemical Co. and used as received. (μ-S)2Fe2(CO)650 and μ-(1,2-benzenedithiolato)Fe2(CO)651 were prepared as reported previously. (50) Brandt, P. F.; Lesch, D. A.; Stafford, P. R.; Rauchfuss, T. B. Inorg. Synth. 1997, 31, 112–116. (51) Cabeza, J. A.; Martinez-Garcia, M. A.; Riera, V.; Ardura, D.; Garcia-Granda, S. Organometallics 1998, 17, 1471–1477.

Article 1

H and 13C NMR spectra were obtained with a Bruker DRX-500 NMR spectrometer with chemical shifts referenced to TMS. IR spectra were collected on a Nicolet 380 FT-IR in heptane solution between two KBr salt plates. High-resolution mass spectra were collected on a JEOL HX110 EB Sector mass spectrometer using direct insertion. Elemental microanalyses were performed at Columbia Analytical Services, Tucson, AZ. Typical Procedure for Base-Induced Addition of (μ-SH)2Fe2(CO)6, (7) to Quinones. Synthesis of 9. A dry, 100 mL Schlenk flask with a stir bar and serum caps was charged with (μ-S)2Fe2(CO)6 (500 mg, 1.45 mmol) and flushed with nitrogen. THF (70 mL) was added by syringe and the solution cooled to -78 C. To this solution was added 1.0 M LiBEt3H in THF (3.0 mL, 3.0 mmol), resulting in a red to green color change. After the solution was stirred for 20 min, CF3COOH (0.23 mL, 3.0 mmol) was added by syringe, which caused a color change to red. This solution was stirred for 10 min at -78 C; then 1,4-naphthoquinone (920 mg 5.82 mmol) and piperidine (0.1 mL, 1.0 mmol) were added to the reaction mixture successively. The reaction mixture was stirred for 30 min at -78 C and then was stirred and warmed to room temperature for 48 h. Subsequently, the reaction mixture was evaporated at reduced pressure to give a dark oil, which was purified by column chromatography on silica gel. Dichloromethane eluted a dark yellow band which on concentration gave [μ-2,3-(naphthalene-1,4-diol)dithiolato]bis(tricarbonyliron) (9) as a brown solid (556 mg, 76% yield). This compound was recrystallized from dichloromethane and hexane. [μ-2,3-(Benzene-1,4-diol)dithiolato]bis(tricarbonyliron) (3a): dark red crystals recrystallized from dichloromethane and hexanes; 1H NMR (CDCl3, 23 C, 500 MHz) δ 6.21 (s, 1 H), 4.84 (s, 1 H); 13C NMR (CDCl3, 23 C, 125 MHz) δ 207.0 (s, Cq, Fe(CO)3), 148.3 (s, Cq, C-OH), 130.4 (s, Cq, C-SFe), 118.2 (s, CH); IR (KBr; cm-1) 2080 (100), 2043 (100), 2013 (100), 1982 (100) (ν(CtO)). [μ-2,3-(5-Methoxybenzene-1,4-diol)dithiolato]bis(tricarbonyliron) (3b): brown powder; 1H NMR (CDCl3, 23 C, 500 MHz) δ 5.79 (s, 1 H), 5.52 and 4.86 (s each, OH, 1:1 H), 3.75 (OCH3, 3 H); 13C NMR (CDCl3, 23 C, 125 MHz) δ 207.3 (s, Cq, Fe(CO)3), 147.4 and 147.1 (s each, Cq each, C-OH), 139.5 (s, Cq, C-OMe), 129.2 and 120.0 (s each, Cq each, C-SFe), 98.9 (s, CH), 56.4 (s, Cq, OCH3); IR (KBr; cm-1) 2074 (100), 2036 (100), 1995 (100) (ν(CtO)); mass spectrum (EI) m/z (relative intensity) 482 (Mþ, 5), 454 (Mþ - CO, 40), 426 (Mþ - 2CO, 30), 398 (Mþ - 3CO, 20), 370 (Mþ - 4CO, 32), 342 (Mþ - 5CO, 40), 314 (Mþ - 6CO, 100), 299 (18), 225 (10), 176 (Fe2S2þ, 20), 144 (Fe2Sþ, 5), 56 (Feþ, 8); HRMS (DIP-EI) m/z M - Hþ, C13H6Fe2O9S2 calcd 481.8153, found 481.8134. [μ-2,3-(5-Methylbenzene-1,4-diol)dithiolato]bis(tricarbonyliron) (3c): orange crystals recrystallized from dichloromethane and pentane; 1H NMR (CDCl3, 23 C, 500 MHz) δ 6.06 (s, 1 H), 4.82 and 4.72 (d each, OH, 1:1 H), 1.97 (CH3, 3 H); 13C NMR (CDCl3, 23 C, 125 MHz) δ 207.1 (s, Cq, Fe(CO)3), 147.5 and 147.0 (s each, Cq each, C-OH), 129.6 and 128.3 (s each, Cq each, C-SFe), 126.6 (s, Cq, C-Me), 118.9 (s, CH), 15.8 (CH3); IR (KBr; cm-1) 2078 (100), 2042 (100), 2009 (100) (ν(CtO)); mass spectrum (EI) m/z (relative intensity) 466 (Mþ, 5), 438 (Mþ CO, 45), 410 (Mþ - 2CO, 33), 382 (Mþ - 3CO, 22), 354 (Mþ 4CO, 30), 326 (Mþ - 5CO, 55), 298 (Mþ - 6CO, 100), 252 (15), 242 (10), 176 (Fe2S2þ, 45), 144 (Fe2Sþ, 15), 56 (Feþ, 10); HRMS (DIP-EI) m/z M - Hþ, C13H6Fe2O8S2 calcd 465.8203, found 465.8191. [μ-2,3-(5-tert-Butylbenzene-1,4-diol)dithiolato]bis(tricarbonyliron) (3d): yellow powder; 1H NMR (CDCl3, 23 C, 500 MHz) δ 6.18 (s, 1 H), 5.06 and 4.72 (s each, OH, 1:1 H), 1.25 (tBu, 9 H); 13C NMR (CDCl3, 23 C, 125 MHz) δ 207.1 (s, Cq, Fe(CO)3), 147.5 and 147.2 (s each, Cq each, C-OH), 140.5 (s, Cq, C-tBu), 130.7 and 125.7 (s each, Cq each, C-SFe), 116.1 (s, CH), 35.1 (s, Cq, C-(CH3)3), 29.2 (CH3); IR (KBr; cm-1) 2081 (100), 2048 (100), 1997 (100) (ν(CtO)). Anal. Calcd for C16H12Fe2O8S2: C, 37.82; H, 2.38. Found: C, 38.07; H, 2.63.


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[μ-2,3-(5-Chlorobenzene-1,4-diol)dithiolato]bis(tricarbonyliron) (3e): dark red crystals (single crystals were obtained by recrystallization from dichloromethane and hexanes); 1H NMR (CDCl3, 23 C, 500 MHz) δ 6.35 (s, 1 H), 5.57 and 4.87 (s each, OH, 1:1 H); 13C NMR (CDCl3, 23 C, 125 MHz) δ 206.8 (s, Cq, Fe(CO)3), 147.3 and 144.3 (s each, Cq each, C-OH), 132.8 (s, Cq, C-Cl), 131.4 and 120.9 (s each, Cq each, C-SFe), 116.9 (s, CH); IR (KBr; cm-1) 2082 (100), 2043 (100), 1993 (100) (ν(CtO)). Anal. Calcd for C12H3ClFe2O8S2: C, 29.63; H, 0.62. Found: C, 29.68; H, 0.87. [μ-2,3-(5-Bromobenzene-1,4-diol)dithiolato]bis(tricarbonyliron) (3f): dark red crystals, recrystallized from dichloromethane and hexanes. 1H NMR (CDCl3, 23 C, 500 MHz) δ 6.49 (s, 1 H), 5.34 and 4.84 (s each, OH, 1:1 H); 13C NMR (CDCl3, 23 C, 125 MHz) δ 206.8 (s, Cq, Fe(CO)3), 147.7 and 145.1 (s each, Cq each, C-OH), 132.4 and 132.3 (s each, Cq each, C-SFe), 119.9 (s, CH), 110.6 (s, Cq, C-Cl); IR (KBr; cm-1) 2082 (100), 2049 (100), 2021 (100) (ν(CtO)). Anal. Calcd for C12H3BrFe2O8S2: C, 27.15; H, 0.57. Found: C, 27.02; H, 0.77. [μ-2,3-(Naphthalene-1,4-diol)dithiolato]bis(tricarbonyliron) (9): dark brown crystals, recrystallized from dichloromethane and hexanes; 1H NMR (CDCl3, 23 C, 500 MHz) δ 7.86 and 7.44 (dd each, 1:1 H), 5.44 (s, 1 H); 13C NMR (CDCl3, 23 C, 125 MHz) δ 207.2 (s, Cq, Fe(CO)3), 144.7 (s, Cq, C-OH), 128.1 and 122.6 (s, CH), 125.9 (s, Cq), 119.1 (s, Cq, C-SFe); IR (KBr; cm-1) 2080 (71), 2034 (89), 1985 (90) (ν(CtO)); mass spectrum (EI) m/z (relative intensity) 502 (Mþ, 10), 474 (Mþ - CO, 11), 446 (Mþ 2CO, 20), 418 (Mþ - 3CO, 10), 390 (Mþ - 4CO, 16), 362 (Mþ 5CO, 16), 334 (Mþ - 6CO, 10), 288 (8), 244 (16), 176 (Fe2S2þ, 24), 144 (Fe2Sþ, 4), 56 (Feþ, 1); HRMS (DIP-EI) m/z M - Hþ, C16H6Fe2O8S2 calcd 501.8203, found 501.8210. Typical Procedure for Photoinduced Addition of (μ-S)2Fe2(CO)6 to Quinones. [μ-2,3-(1,4-Dimethoxybenzene)dithiolato]bis(tricarbonyliron) (2b, Red Powder). (μ-S2)Fe2(CO)6 (5; 100 mg, 0.29 mmol) and 1,4-benzoquinone (4a; 63 mg, 0.58 mmol) were dissolved in THF (10 mL) in a 25 mL quartz tube. The resulting solution was irradiated for 3 h at room temperature in a Srinivasan-Griffith Rayonet photochemical reactor equipped with 350 nm lamps. The volatiles were removed under vacuum, and K2CO3 (80 mg, 0.58 mmol), iodomethane (0.2 mL, 3.21 mmol), and 10 mL of methanol were added. The solution was stirred at room temperature overnight. The volatiles were removed under vacuum, and the residue was separated by column chromatography on silica gel with hexanes as eluent to give the red powder 2b (26 mg, 19% yield). 1H NMR (CDCl3, 23 C, 500 MHz): δ 6.28 (s, 1 H), 3.74 (s, 1 H). 13C NMR (CDCl3, 23 C, 125 MHz): δ 207.5 (s, Cq, Fe(CO)3), 152.2 (s, Cq, C-OH), 135.2 (s, Cq, C-SFe), 112.7 (s, CH), 56.4 (s, CH3). IR (CHCl; cm-1): 2078 (100), 2043 (100), 2005 (100) (ν(CtO)). Anal. Calcd for C12H3ClFe2O8S2: C, 35.02; H, 1.68. Found: C, 35.11; H, 2.19. [μ-2,3-(Naphthalene-1,4-diol)dithiolato]bis(tricarbonyliron) (9). (μ-S2)Fe2(CO)6 (100 mg, 0.29 mmol) and 1,4-naphthoquinone (92 mg, 0.58 mmol) were dissolved in THF (100 mL) in a 200 mL Schlenk flask. The resulting solution was irradiated for 3.0 h at room temperature in a Srinivasan-Griffith Rayonet photochemical reactor equipped with 350 nm lamps. The volatiles were removed under vacuum, and the residue was separated by column over silica gel by using CH2Cl2 solvent as eluent to give 9 (63 mg, 43% yield). Electrochemistry. For electrochemical experiments, the source and treatment of the solvent and supporting electrolyte have been described earlier.31 Experiments were conducted in acetonitrile or dichloromethane with 0.10 M tetrabutylammonium hexafluorophosphate (Bu4NPF6) as supporting electrolyte. Electrodes, cells, instrumentation, and electrochemical procedures have been described.52 In the present study the potentiostat was an EG&G PAR Model 2273. The working (52) Macias-Ruvalcaba, N. A.; Evans, D. H. J. Phys. Chem. B 2005, 109, 14642–14647.

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electrode was a glassy-carbon electrode (0.071 cm2) or a mercury film on a gold-disk electrode (0.080 cm2), prepared as described earlier.31 Its area was ascertained by studies of the oxidation of ferrocene, whose diffusion coefficient is known.52 Evaluation of solution resistance was carried out as described earlier,52 and the resistance was partially compensated by electronic resistance compensation, with the remainder of the resistance applied when simulating the data. The laboratory reference electrode was a silver wire in contact with 0.010 M AgNO3, 0.10 M NBu4PF6, acetonitrile (AgRE). Its potential was frequently measured with respect to the reversible ferrocene/ferrocenium couple, in acetonitrile or dichloromethane, and all potentials are referred to ferrocene. Voltammetric experiments were carried out at room temperature, except for those in dichloromethane with added acid, for which the temperature was 298 K. Digital simulations were conducted with DigiElch, version 3.0, a software package for the digital simulation of common electrochemical experiments (http://www.digielch.de).53 The fitting routine in that program was used to establish the final best-fit parameter values for many of the variables. Photoelectron Spectroscopy. All samples sublimed cleanly, with no visible changes in the spectra during data collection (53) Rudolph, M. J. J. Electroanal. Chem. 2003, 543, 23–29. (54) Siegbahn, K.; Nordling, C.; Fahlman, A.; Nordberg, R.; Hamrin, K.; Hedman, J.; Johansson, G.; Bergmark, T.; Karlsson, S. E.; Lindgren, I.; Lindberg, B. Nova Acta Regiae Soc. Sci. Ups. 1967, 20, 282. (55) Lichtenberger, D. L.; Kellogg, G. E.; Kristofzski, J. G.; Page, D.; Turner, S.; Klinger, G.; Lorenzen, J. Rev. Sci. Instrum. 1986, 57, 2366. (56) Cranswick, M. A.; Dawson, A.; Cooney, J. J. A.; Gruhn, N. E.; Lichtenberger, D. L.; Enemark, J. H. Inorg. Chem. 2007, 46, 10639– 10646.

Chen et al. after initial observation of ionizations from the diiron complex. Photoelectron spectra were recorded using an instrument that features a 36 cm hemispherical analyzer (McPherson)54 with custom-designed photon source, sample cells, detection and control electronics, calibration, and data analysis, as described previously.55,56 The solid line is the best fit of the He I data (shown as vertical dashes) with a minimum number of Gaussian peaks to represent the contour of the ionization intensity. The dashed line is a similar representation of the contour of ionization intensity obtained with the He II photon source, scaled to match the low ionization energy intensity in the He I spectrum for visual comparison of the change in relative intensity at higher ionization energies. Computational Methods. All computational methods have been described previously.13 See the Supporting Information for the specific computational details.

Acknowledgment. The support of the National Science Foundation through the Collaborative Research in Chemistry program, Grant No. CHE 0527003, is gratefully acknowledged. Supporting Information Available: Text, figures, and tables giving further details on the computational methodology and the photoelectron spectroscopy experiments, Cartesian coordinates and electronic energies of all reported calculated structures, and IR spectra of 2a,b in acetonitrile with a large excess of added acetic acid. This material is available free of charge via the Internet at http://pubs.acs.org.

Synthesis of Diiron Hydrogenase Mimics Bearing Hydroquinone and

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