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Electrochemical, Spectroscopic, and Computational Study of Bis(μmethylthiolato)diironhexacarbonyl: Homoassociative Stabilization of the Dianion and a Chemically Reversible Reduction/Reoxidation Cycle Orrasa In-noi,† Kenneth J. Haller,*,† Gabriel B. Hall,‡ William P. Brezinski,‡ Jacob M. Marx,‡ Taka Sakamoto,‡ Dennis H. Evans,§ Richard S. Glass,‡ and Dennis L. Lichtenberger*,‡ †

School of Chemistry, Institute of Science, Suranaree University of Technology, Nakhon Ratchasima 30000 Thailand Department of Chemistry and Biochemistry, The University of Arizona, Tucson, Arizona 85721, United States § Department of Chemistry, Purdue University, West Lafayette, Indiana 47907, United States ‡

S Supporting Information *

ABSTRACT: The redox characteristics of (μ-SMe)2Fe2(CO)6 from the 1+ to 2− charge states are reported. This [2Fe-2S] compound is related to the active sites of [FeFe]-hydrogenases but notably without a linker between the sulfur atoms. The 1+ charge state was studied both by ionization in the gas phase by photoelectron spectroscopy and by oxidation in the solution phase by cyclic voltammetry. The adiabatic ionization is to a cation whose structure features a semibridging carbonyl, similar to the structure of the active site of [FeFe]-hydrogenases in the same oxidation state. The reduction of the compound by cyclic voltammetry gives an electrochemically irreversible cathodic peak, which often suggests disproportionation or other irreversible chemical processes in this class of molecules. However, the return scan through electrochemically irreversible oxidation peaks that occur at potentials around 1 V more positive than the reduction led to the recovery of the initial neutral compound. The dependence of the CVs on scan rate, IR spectroelectrochemistry of reduction and oxidation cycles, chronocoulometry, and DFT computations indicate a mechanism in which stabilization of the dianion plays a key role. Initial one-electron reduction of the compound is accompanied in the same cathodic peak with a second slower electron reduction to the dianion. Geometric reorganization and solvation stabilize the [2Fe2S]2− dianion such that the potential for addition of the second electron is slightly less negative than that of the first (potential inversion). The return oxidation peaks at more positive potentials follow from rapid pairing of the dianion with another neutral molecule in solution (termed homoassociation) to form a stabilized [4Fe-4S]2− dianion. Two one-electron oxidations of this [4Fe-4S]2− dianion result in regeneration of the initial neutral compound. The implications of this homoassociation for the [FeFe]-hydrogenase enzyme, in which the H-cluster active site features a [2Fe-2S] site associated with a [4Fe-4S] cubane cluster via a thiolate bridge, are discussed.

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organometallic [2Fe-2S] clusters.1,2,6,7,15 The oxidation and reduction properties of these species are central to their function. The majority of the mimics that have been studied to this time have a linker between the two sulfur atoms of the [2Fe2S] core, as depicted by the dashed line in Chart 1.16 This linker can have both a direct functional role and a structural role in the reduction chemistry. The most notable functional role is incorporation of an amine in the linker that can act as a protonation site and proton relay to the [2Fe-2S] core.17,18 The

INTRODUCTION The [2Fe-2S] cluster unit has important roles in biological chemistry, and the [2Fe-2S] unit also has a richly developed organometallic chemistry. Most relevant to the present study are the [2Fe-2S] clusters in the active sites of [FeFe]hydrogenase enzymes that catalyze the reversible reduction of protons to hydrogen.1−7 In organometallic chemistry, the structure of the [2Fe-2S] cluster, μ-S2Fe2(CO)6 (1, Chart 1) was first published in 1965,8−10 and some of the early chemistry of this class of complexes of the general form shown in Chart 1 was developed by Dietmar Seyferth.11−14 More recently, the structural similarity of this class of complexes to the active site of [FeFe]-hydrogenases has inspired widespread investigations into the electrocatalytic reduction of protons to hydrogen by © 2014 American Chemical Society

Special Issue: Organometallic Electrochemistry Received: April 18, 2014 Published: July 21, 2014 5009

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Reported here are the results of electrochemical rate measurements, electrochemical simulation, and spectroelectrochemical characterization of reduction of the MeMe compound. The results provide a systematic picture of the relative energies and structures of the MeMe compound over a range of four molecule charge states from 1+ to 2−. Interestingly, the model for this chemistry that emerges from these results reconciles the apparently disparate results of previous studies. The key is a homoassociation process in which the dianion rapidly pairs with another neutral molecule to stabilize the negative charge. Most significant is the complete chemical reversibility of the reduction when the reoxidation is carried out by the methods described here.

Chart 1

linker can also play a structural role in stabilizing the catalyst19 or influencing the geometries of intermediates and transition states, such as favoring rearrangement to the “rotated” structure of the [FeFe]-hydrogenase active site shown in Chart 1 with a bridging carbonyl ligand.20 What are the consequences of removing the linker between the sulfur atoms? Does the extra degree of flexibility for the [2Fe-2S] cluster favor the kinetics for catalysis, or does the lack of a linker simply destabilize the cluster toward disproportionation? The archetype alkylthiolato molecule for probing these questions is (μ-SMe)2Fe2(CO)6 (MeMe, Chart 2). The

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RESULTS The electrochemical observations will be discussed in parallel with the results of the computations to illustrate the interplay between experiment and theory in developing a more complete understanding of the redox behavior. It is helpful first to review the known structure, relative isomer energies, and spectroscopic properties of the MeMe compound in relation to the computations. The ability of the computations to account for the structure and carbonyl stretching frequencies contributes to interpretation of the IR spectroelectrochemistry, and the gas phase ionization energies provide a well-defined measure of the change in energy with change in electron configuration and geometry. Knowledge of the relationship between these experimental and computational energies is essential for understanding the standard oxidation and reduction potentials in solution. Thermodynamics of the Neutral Molecule. The MeMe compound was prepared in high yield by modification of a literature procedure as shown in Scheme 1.11 The axial vs

Chart 2

Scheme 1a

foundations for the present investigation have been established by several previous studies of the chemical and physical properties of this molecule.21−30 This compound is particularly attractive for the present study because it is closely related to the much-studied compound (μ-propane-1,3-dithiolato)Fe2(CO)6 (PDT, Chart 2). Both the MeMe and PDT compounds have alkyls bound to the sulfur atoms, but only the PDT compound links the sulfur atoms. The extra degree of freedom of the MeMe compound is evidenced by the existence of additional isomers as shown in Chart 2. Neither of the characterized isomers has a diaxial substitution as shown by PDT. An early electrochemical study reported a one-electron reduction of the MeMe compound and that 100% of the “starting material was recovered after exhaustive controlled potential reduction followed by oxidation.”30 This is an interesting result considering that cyclic voltammety of PDT shows a reduction that is quasireversible even at fast scan rates, and chemical reduction under CO transforms the compound to a [4Fe-4S] species that has been structurally characterized.31 Later, electrochemical studies of the MeMe compound reported that it undergoes a two-electron reduction and that the starting material was not recovered on reoxidation in complete contrast to the previous report.29,32 A study of the related bis-ethanethiolato molecule (μ-SEt)2Fe2(CO)6 observed that the first reduction peak was more irreversible than that of PDT, but about 75% of the starting material was recovered following a reoxidation cycle at more positive potentials.33

(a) THF, LiEt3BH, −78 °C, 30 min; (b) CH3I, −78 °C, 20 min; ambient temperature, 3 h. The axial vs equatorial position of the Me groups is not specified in the Scheme.

a

equatorial substitution of the methyl groups on the sulfur atoms (notation defined in Chart 2) is not specified in the scheme. The ratio of the (aMe)(eMe) to (eMe)(eMe) isomers obtained from separate preparations were determined by 1H NMR spectroscopy and ranged from 7:1 to 20:1 in this work. Previous preparations by a variety of methods also find the (aMe)(eMe) isomer to be dominant, with ratios ranging from 2.7:1 to 12:1.11,12,22 The structure with both alkyl groups in axial positions has not been observed and apparently is not favored. Consistent with this behavior, the (aMe)(eMe) isomer is calculated to have a Gibbs energy just 2 kJ/mol more stable than the (eMe)(eMe) isomer, which would give a thermodynamic product ratio of 2:1 at room temperature. The (aMe)(aMe) isomer is calculated to be 31 kJ/mol less stable and therefore not observed at ambient temperature. The implication is that initial axial substitution is slightly preferred, 5010

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forces do not have a major influence on the structural parameters. For later evaluation of the oxidation and reduction potentials, the optimizations were carried out in both gas phase and solvent as described in the Experimental Section, and it was found that the distances changed by less than 0.002 Å from the gas phase to solution. The carbonyl stretching frequencies give an indication of the electron richness at the metal center, and the splittings and relative intensities of the carbonyl absorptions in the IR spectrum are sensitive to the angular geometries and couplings between the vibrations. In order to model the carbonyl stretching region of the IR spectrum, computations must be able to account for the electron distribution, the ground state geometry, and the curvature of the potential energy surface in the region of the equilibrium geometry. Evaluation of the carbonyl stretching region in the IR spectra of the neutral molecules provides a framework for evaluation of the IR spectroelectrochemistry results on reduction to the negative ions discussed later. The IR spectrum in the carbonyl stretching region for the (aMe)(eMe) compound is shown in Figure 1, along with the modeling from the computations. The computations reproduce the frequencies within 0.3% and match the splitting and intensity pattern well.

probably both thermodynamically and kinetically, but diaxial substitution is disfavored by steric repulsion. It should be noted that the propanedithiolato linker between the sulfur atoms in the PDT compound constrains the conformation to the diaxial substitution (Chart 2). The relative energies of key isomers of the MeMe molecule from the cation to the dianion are collected in Table S1 in the Supporting Information. A previous NMR study has shown that the MeMe compound does not undergo axial−equatorial ligand exchange at temperatures up to 80 °C, at which temperature the compound decomposes.25 The activation energy for transition from the (aMe)(eMe) isomer to the (eMe)(eMe) isomer is calculated to be 112 kJ/mol. The calculated lowest energy path between the isomers is not the breaking of one Fe−S bond followed by rotation of the SMe and then reconnection but is instead a simple inversion of the trigonal sulfur atom through a trigonal planar transition state (the sulfur atom, two iron atoms, and methyl carbon atom all in a plane) similar to the inversion of ammonia. Because of the high energy of this transition state, the isomers can be separated by chromatography and studied individually. In addition to the ability of the computations to account for the relative stabilities of the isomers, the computations also compare well with the experimental structure. Interestingly, the structure of the MeMe compound that has been reported34 shows the (eMe)(eMe) isomer, which is the minor of the two isomers from all preparations. Some key structural parameters are shown in Table 1 and compared with those of PDT. The Table 1. Selected Structural Featuresa (eMe)(eMe)b Fe−Fe Fe−S S···S Fe−Ca Fe−Cb C−O S−Fe−S Fe−Fe−Ca

PDTc

X-ray

DFT

X-ray

2.52 2.26 2.76 1.81 1.78 1.14 75 150

2.51 2.28 2.77 1.78 1.76 1.16 75 151

2.51 2.25 3.05 1.80 1.80 1.14 85 148

Figure 1. IR spectrum in the carbonyl region of the (aMe)(eMe) isomer of (μ-SMe)2Fe2(CO)6. The black lines are the experimental absorptions (in hexane), and the light blue lines are the DFTcalculated absorptions (gas phase, scaled by 0.3%, arbitrary width).

Characterization of the Positive Ion by Gas-Phase Photoelectron Spectroscopy. Photoelectron spectroscopy is a powerful probe of the electronic structure of a molecule, and the results are commonly discussed in terms of the characters of the molecular orbitals. However, more strictly speaking, the ionization energies provide the relative energies of the positive ion states, and the ionization peak shapes provide information on the geometric and vibrational relationships of the positive ion states to the neutral molecule. Furthermore, the ionizations provide a well-defined benchmark for the computational electronic energies in the gas phase. The photoelectron spectrum of the MeMe compound is shown in Figure 2. The general atomic orbital characters of the ionizations are indicated by the changes in relative peak intensities between the He I and He II spectra. The ionization cross-section for sulfur-based ionizations decreases dramatically from He I to He II excitation compared to metal carbonyl based ionizations.38−40 The decrease in cross-section for carbon-based ionizations is intermediate to these two. Consequently, as in previous photoelectron studies of related molecules,41−45 the first ionization band spread from about 7 to 9 eV is assigned to ionization predominantly from the seven iron-based orbitals (combination of three orbitals from the formal d6 electron configurations of each Fe center plus the metal−metal bond). The sulfur-based ionizations from about 9.5 to 11 eV show a dramatic decrease in relative intensity from

a

Distances are in Angstroms and angles in degrees. Ca are the apical carbonyls, and Cb are the basal carbonyls. All parameters are averaged over independent molecules in the unit cells for similar bond types in the molecules. bRef 34. cRef 35.

structure is for the most part typical of this class of butterfly complexes. For instance, the average Fe−Fe and Fe−S distances of 154 compounds of the general form of compounds shown in Chart 1 with L = CO are 2.51 and 2.26 Å, respectively,36,37 which are the same as those found for these compounds within 0.01 Å. The primary difference in the structure of the MeMe compound is the significantly shorter nonbonded S···S distance of 2.76 Å. The average S···S distance of the 58 compounds in the database with a three-carbon linker is 3.02 Å and with a two-carbon linker is 2.92 Å. The shorter S···S distances in the MeMe compound implies that some strain has been relieved. The (eR)(eR) isomers generally have the shortest S···S distances. The agreement between the density functional theory (DFT) geometry optimizations and the crystal structures, including the nonbonded S···S distances, is achieved with geometries that are optimized in the gas phase, suggesting that crystal packing 5011

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the sulfur atom with an equatorial substituent that serves to destabilize the HOMO. This mixing is not present when the substituents are axial, as in the case of PDT. A point to note for later discussion is that this orbital is spatially accessible and energetically favorable for intermolecular interactions. Most important for understanding the oxidation potential is the nature of adiabatic ionization. Because of the metal−metal bonding nature of the HOMO of these compounds, removing an electron from HOMO is expected to substantially weaken the metal−metal bond and result in a geometric reorganization that lengthens the Fe−Fe distance. However, the computations find that the most stable reorganization results in the “rotation” of an Fe(CO)3 group and shift of a terminal carbonyl to a semibridging position, where it is able to help compensate for the electron deficiency created at an iron center with oxidation. This reorganization was found computationally to be spontaneous without a barrier starting from several of the neutral molecule isomers. This cation structure for the (aMe)(eMe) compound is shown in Figure 4.

Figure 2. He I (blue) and He II (red) photoelectron spectra of the (aMe)(eMe) isomer of (μ-SMe)2Fe2(CO)6. The arrows point to the DFT-calculated vertical ionization energies (VIE) and adiabatic ionization energies (AIE). The vertical lines at the top of the spectra show the pattern of Kohn−Sham orbital energies approximately aligned with the spectra.

He I to He II excitation. Overall, the changes in relative ionization intensities from He I to He II excitation are not as drastic as expected for pure atomic orbitals, indicating appreciable delocalization of the orbitals throughout the molecule. Figure 2 also shows the calculated vertical ionization energies (VIE), adiabatic ionization energies (AIE), and the pattern of the Kohn−Sham orbital energies from the DFT computations. A molecular orbital energy level diagram of this compound has been published previously.26 Koopmans’ theorem relating Hartree−Fock orbital energies to ionization energies does not strictly apply to Kohn−Sham orbital energies, and there are further issues of electron relaxation and correlation that obviate a direct comparison to the experiment. Nonetheless, the pattern of the orbital energies and the orbital characters often provide insight into the ionizations. Figure 2 shows that the orbital energies reproduce the separation of the highest occupied molecular orbital (HOMO) ionization to low energy and that the cluster of six iron-based ionizations are followed by sulfurbased ionizations. Isosurface plots of these orbitals and the calculated energies are shown in Supporting Information, Figure S1. A significant feature is the shoulder on the low ionization energy side of the band, which represents ionization from an orbital that is predominantly the metal−metal bond. This shoulder is more prominent than generally observed in the spectra of related compounds that have a linker between the sulfur atoms. A plot of the HOMO, shown in Figure 3, illustrates the delocalization of the predominantly metal−metal bond, with a substantial contribution from the filled p orbital of

Figure 4. Optimized calculated cation structure of [(μSMe)2Fe2(CO)6]1+.

Figure 5. Cyclic voltammograms for oxidation (left) and reduction (right) of (μ-SMe)2Fe2(CO)6 with the DFT-calculated standard potentials indicated by the arrows. Glassy carbon working electrode, scan rate 100 mV/s, 1 mM sample, 0.10 M n-Bu4NPF6/acetonitrile, 298 K, and background corrected.

Solution Oxidation. Figure 5 shows the cyclic voltammogram of the MeMe compound. Solution oxidation potentials are related to gas-phase adiabatic ionization energies by the differential effects of solvation on the neutral and ion species and by the thermodynamic factors that contribute to the Gibbs energies. The greater solvation energy of the positive ion compared to that of the neutral molecule is calculated to lower the oxidation potential by 1.92 V in acetonitrile compared to

Figure 3. Highest occupied (HOMO) and lowest unoccupied (LUMO) molecular orbital plots of (μ-SMe)2Fe2(CO)6. The isosurface values are ±0.04. 5012

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that the dissociation energy is on the same scale as the axial/ equatorial exchange energy, which as mentioned before is not observed up to 80 °C. Another consideration is the possibility of carbonyl dissociation with reduction giving an irreversible process. Cyclic voltammograms for reduction of the PDT molecule under a CO atmosphere show increased reversibility, consistent with this hypothesis.51 Capon et al. calculated an energy difference for loss of a CO molecule from the PDT anion of 77 kJ/mol uphill.52 The computational model used here similarly calculates an energy difference of 79 kJ/mol. Furthermore, a linear transit computation of the CO dissociation yields an activation Gibbs energy of 101 kJ/mol:

the gas phase ionization energy. Further adjustments due to optimization of the geometries in solution, differences in zeropoint vibrational energies, and thermal enthalpy and entropy contributions at standard temperature and pressure add up to less than 0.01 V for oxidation of the MeMe compound. Nonetheless, these contributions were included in all calculations of oxidation and reduction potentials in this study. Our experience has been that potentials calculated by this model agree well with the location of electrochemical processes in the cyclic voltammetry (CV) taken in acetonitrile, as shown in Figure 5, without further consideration of nonideal solvent effects on the activity coefficients of the species. Figure 5 shows that the calculated oxidation potential for the MeMe compound is approximately at the onset of the oxidation current. The oxidation is fully irreversible, and a more detailed exploration of the oxidation process has not been pursued. Solution Reductions. The reduction behavior is of primary interest for the electrocatalytic reduction of protons to hydrogen. The cyclic voltammogram for reduction of the MeMe compound is shown on the right side of Figure 5. Mathieu et al. reported that the number of electrons involved in this cathodic peak is two.29 The present results agree with two electrons based on the comparison of the peak height of this reduction relative to that of ferrocene oxidation and taking into account the relative diffusion coefficients in acetonitrile (2.5 × 10−5 cm2/s for ferrocene,46 1.59 × 10−5 cm2/s for (1,2ethanedithiolato)Fe2(CO)647) and the different electron transfer rates (vida infra). The reduction is electrochemically irreversible with no anodic peaks in the region of −1.5 to −0.8 V at this scan rate and with new oxidation peaks appearing at potentials more positive than −0.8 V. These observations indicate that chemical change with reduction is taking place that significantly stabilizes the dianion charge. The early concern was that the [(μ-SMe)2Fe2(CO)6]2− dianion produced by two-electron reduction was splitting into two [(SMe)Fe(CO)3]− monomer anions, similar to the known splitting of [Cp 2 Fe 2 (CO) 4 ] 2− into two [CpFe(CO) 2 ] − monomer anions.48 The [(SMe)Fe(CO)3]− monomer anion is an 18 e− complex if the thiolato is considered a four electron donor, and if a solvent acetonitrile molecule occupies a coordination site, the description is that of a familiar d8 fivecoordinate complex. This thought has further credibility considering the related [(SMe)Fe(CO)4]− anion has been prepared by Darensbourg but by a much different route.49 However, formation of the monomer anions by reduction of the MeMe compound is not supported by computations. The following two dissociations were modeled in acetonitrile:

[(SMe)2 Fe2(CO)6 ]− ⇌ [(SMe)2 Fe 2(CO)5 ]− + CO ΔG⧧ = 101 kJ/mol

These energies are again too high for CO dissociation to contribute significantly to the cyclic voltammograms at ambient temperature. Interestingly, the addition of CO provides slight stabilization of the anion charge: [(SMe)2 Fe2(CO)6 ]− + CO ⇌ [(SMe)2 Fe 2(CO)7 ]− ΔG⧧ = −6 kJ/mol

This stabilization of the anion under CO can contribute to the increased reversibility in the cyclic voltammogram. Further discussion of reduction chemistry in the presence of CO or other small molecules is beyond the scope of this report, but it is noted that associative interactions can stabilize the negative charge. The main point is that the electronic structure and bonding of this MeMe molecule does not favor disproportionation into monomers or carbonyl dissociation with reduction. Experimental evidence in support of this point follows. Reduction/Return Oxidation/Reduction Cycle. While return oxidation in the vicinity of the reduction is not observed in Figure 5, the return oxidation peaks at ∼ −0.75 and −0.5 V after reduction play an important role in recovery of the neutral species. The first hint of a chemically reversible cycle came from performing a 3-segment cyclic voltammetry experiment, as displayed in Figure 6. The first segment of the CV scan runs through the reduction, the second segment scans back through the oxidations at ∼ −0.75 and −0.5 V, and a third segment then scans back through the reduction. Most of the reductive current is recovered on the third segment. However, if the second segment is stopped short of the return oxidations at −0.75 and −0.5 V, then observable loss of cathodic current occurs on the scan back through the reduction. The contribution of diffusion to the current was equalized by pausing at −1.5 V for the same six seconds taken in the full scan from this point and back. At least a 30 s delay was necessary for diffusion to achieve the same MeMe concentration as scanning through the −0.75 and −0.5 V peaks. As an additional check, a simulation was constructed in which the return oxidations at −0.75 and −0.5 V were either completely reversible in returning the MeMe compound or completely irreversible. The simulation current in the third segment for the completely chemically reversible situation closely matches the experiment. The simulation current in the third segment for the irreversible situation, due entirely to diffusion, is less than that in the experiment. Thus, the MeMe compound is recycled after reduction on the CV time scale by

[(μ‐SMe)2 Fe2(CO)6 ]2 − ⇌ 2[(SMe)Fe(CO)3 ]−

[(μ‐SMe)2 Fe2(CO)6 ]2 − + 2NCMe ⇌ 2[(SMe)Fe(CO)3 NCMe]−

In both cases, dissociation is disfavored computationally by 120 kJ/mol. The largest uncertainty in modeling dissociative or associative reactions in solution is the handling of the solution translational and rotational entropy contributions to the Gibbs energy.50 In the first reaction, translational entropy favors the product, and in the second reaction, translational entropy favors the reactants. The result that the same Gibbs energy change is obtained for both reactions indicates that solvation and the method we developed previously for the solution entropy50 are performing well for these complexes. There is little uncertainty 5013

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Scheme 2. Calculated Structures and Potentials for the Reduction of (μ-SMe)2Fe2(CO)6a

a

Figure 6. Three-segment scans of (μ-SMe)2Fe2(CO)6 with a scan rate 0.5 V/s. The full sweep (0 → −2 → 0 → −2 V) is the blue line with the third segment in green. In the experiment, the red line is a similar 3-segment sweep but with a pause at −1.5 V for 6 s to equal the scan time of the full sweep before scanning back negative in the third segment (0 → −2 → −1.5 (6 s) → −2 V). In the simulation, the full sweep is a fully chemically reversible mechanism that recycles the starting compound through the first and second segments (see text), and the red line is a mechanism in which the return oxidations are fully irreversible and do not recycle the starting compound.

Notice the potential inversion.

reduction peak corresponds to two electrons. At the scan rate of 0.5 V/s shown in Figure 6, a weak return oxidation peak is observed at −1.3 V that is not observed at the slower 0.1 V/s scan rate in Figure 5. This peak corresponds to oxidation of the [2Fe-2S]2− dianion that is captured at the faster scan rate. The width and shape of this return oxidation peak indicates a slow and asymmetric electron transfer for the anion/dianion couple, consistent with the substantial geometry change between the two structures as shown previously.47 Related dianion conformations with the bridging thiolato ligand as the axial or equatorial isomer and with the terminal thiolato methyl group in various rotational conformations about the Fe−S bond are within 10 kJ/mol of the most stable structure and may contribute to the width of the peak. The structure of the dianion with the elongated Fe−Fe distance and all terminal carbonyls is 43 kJ/mol higher in energy with a potential for the 1−/2− couple at −1.95 V and therefore cannot contribute to the CV scans in the potential windows of Figures 5 and 6. The return oxidation peaks at −0.75 and −0.5 V in Figure 6 require additional consideration for stabilization of the dianion charge. Explicit coordination of an acetonitrile solvent molecule to the dianion shown in Scheme 2 does not add appreciable stabilization of the charge. Likewise, pairing of the [n-Bu4N]+ cation from the electrolyte with the dianion is not appreciable in acetonitrile. Another method for stabilization of the dianion charge is coupling with another MeMe molecule. Indeed, it is known that reduction of PDT (under CO) and (μ-1,2ethanedithiolato)Fe2(CO)6 (EDT) ultimately results in dimers of the diiron complexes.31,54 Association of [2Fe-2S] clusters has been shown to stabilize the dianion and lower the reduction potential.55 This MeMe case is different because the sulfur atoms are not linked to each other, and the electrochemistry is not carried out under CO. Nonetheless, computations can probe the feasibility of homoassociation of the MeMe dianion with a second MeMe molecule. The simplest initial association is for the sulfur atom that is detached from one iron center in the MeMe dianion to coordinate with the iron atom of a neighboring MeMe molecule. This kind of association of the dianion has been implicated in the combined IR spectroelectrochemistry, X-ray absorption fine structure (EXAFS), and

scanning back through the return oxidation peaks. The observation of three return oxidation peaks at the scan rate in Figure 6 (−1.3 V, −0.75 V, and −0.5 V) indicates that the reduced species is transformed to more than one species in which the negative charge is stabilized compared to that in the initial reduction. Anionic Structures and Energies. Computations give insight into the energies and structures of the species probed experimentally by the cyclic voltammograms shown at the top of Figure 6 and provide the basis for the simulation shown at the bottom of Figure 6. The first reduction potential of the MeMe compound is calculated at −1.62 V, as shown in Scheme 2. The most stable anion structure is the (aMe)(eMe) isomer with all terminal carbonyl ligands similar to the neutral molecule but with the Fe−Fe distance lengthened by 0.32 Å (Scheme 2). This structural transformation is consistent with the Fe−Fe antibonding nature of the lowest unoccupied molecular orbital (LUMO) of the neutral molecule shown in Figure 3. The most stable structure of the dianion breaks one Fe−S bond instead and rotates one Fe(CO)3 end of the molecule to place a carbonyl in a fully bridging position (structure shown schematically in Scheme 2). This arrangement serves to stabilize the dianion by reducing the electron−electron repulsion. This structure has been found computationally for other related [2Fe-2S] compounds42,47,52 and characterized crystallographically in the case of (1,2-benzenedithiolato)Fe2(CO)6.53 The stabilization is sufficient in acetonitrile that the calculated potential for addition of the second electron, −1.50 V, is less negative than the potential for addition of the first electron (potential inversion). This potential inversion accounts for the observation that the current in the first 5014

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computational study of EDT.56 For the MeMe molecule, the orbital used for this coordination begins as the HOMO shown in Figure 3, which has significant sulfur lone-pair character according to the computations and is energetically available as shown by the photoelectron spectroscopy. Calculations indicate that the equilibrium constant for this association of the MeMe dianion is greater than 104, as shown in Scheme 3. Unlike the

a [4Fe-4S]2− species. Oxidation of the [4Fe-4S]2− species in two one-electron steps then returns the neutral molecules. The CV depends on the rates of electron transfer, the rate of homoassociation, and the scan rate of the CV. The experimental CVs obtained with scan rates varying over 2 orders of magnitude from 0.1 V/s to 10 V/s are shown in Figure 7 along with a simulation using the mechanism in

Scheme 3. Calculated Gibbs Energy and Structures for the Associated [4Fe-4S]2− Complex Formed from the Neutral and Dianion Molecules of (μ-SMe)2Fe2(CO)6a

The equilibrium constant for dimer formation in acetonitrile is ∼6 × 104.

a

previous examples, each diiron retains six carbonyls in this initial association. The stabilization of the dianion by this association corresponds to a positive shift of greater than 300 mV in the oxidation potential of the dianion. Several related conformations were found within 10 kJ/mol. General Mechanism and Scan Rate Simulation. The favorability of association between the dianion and the neutral molecule suggests the mechanism in Scheme 4 for the

Figure 7. Scan rate comparison of (μ-SMe)2Fe2(CO)6 with current normalized for the scan rate: green, 0.1 V/s; red, 1 V/s; blue 10 V/s. The conditions were the same as those described in Figure 5.

Scheme 4. For the simulation, the potentials for the first and second reductions and the equilibrium constant for homoassociation of the dianion were fixed at the values from the computations (Scheme 4). A fixed diffusion rate of 1.59 × 10−5 cm2/s for all species was taken from a previous rate study of (μ1,2-ethanedithiolato)Fe2(CO)6.47 Further details are provided in Supporting Information. A perfect overlay of the simulation on the experiment is not expected with these constraints, but nonetheless, the simulation captures the major features of the experimental CVs. The position and maximum current of the initial reduction is well matched, as is the return reduction current on the third segment after sweeping through the return oxidations on the second segment (Figure 6). The broad return oxidation peak at −1.3 V observed at fast scan rates is due to the slow electron transfer associated with the distorted [2Fe2S]2− dianion. The −0.75 and −0.5 V return oxidation peaks are due to the first and second oxidations of the homoassociated [4Fe-4S]2− dianion that has formed. On the basis of this simulation, the stabilization of the dianion by homoassociation could be as much as 700 mV. The relative intensities of the [2Fe-2S]2− and [4Fe-4S]2− oxidation peaks are dependent on the homoassociation rate and the scan rate. Note that at the slower scan rate shown in Figure 7 (0.1 V/s), the [2Fe-2S]2− dianion has been completely consumed by homoassociation and that the oxidation peak at −1.3 V is not observed. At the fast scan rate (10 V/s), there is not sufficient time for complete homoassociation, and the oxidation peak at −1.3 V is strong, and the −0.75 and −0.5 peaks are very weak. Intermediate scan rates show a weaker

Scheme 4. Reduction/Oxidation Mechanism for (μSMe)2Fe2(CO)6a

a M is (μ-SMe)2Fe2(CO)6, and M−M is the homoassociated [4Fe-4S] species (perhaps as shown in Scheme 3). The potentials and equilibrium constants are from the computations.

reduction and return oxidation of the MeMe compound. The reduction potentials of the species and equilibrium constants for homoassociation are from the computations. The initial reduction of the molecule to the anion is followed by a second slower reduction due to the geometry reorganization to a stabilized dianion with potential inversion. The dianion is further stabilized by association with another molecule to form 5015

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oxidation peak in the region of −1.5 to −1.3 V that shifts to more positive potential in this region as the scan rate is increased, consistent with the slow electron transfer process for the [2Fe-2S]1−/2− couple. IR Spectroelectrochemistry. The infrared spectrum of the MeMe compound after electrolysis for a period of ∼1 min at a potential ∼0.2 V negative of the cathodic CV peak at −1.6 V is shown in Figure 8. The spectrum is similar to that in a previous

compound, which was also explained by the formation of a similar associated [4Fe-4S] complex.57 However, the dimer of the dianion of the biphenyl complex was interpreted with the loss of one CO ligand, which is not consistent with the chemical reversibility observed in the case of this MeMe compound. At no time was the IR absorption of a free CO molecule in solution observed in this experiment. Controlled Potential Coulometry (CPC). CPC was carried out both from the neutral to the fully reduced species in solution and from the fully reduced species back to the neutral compound, and monitored by IR in the spectroelectrochemical cell. The total charge passed was found to be one electron per molecule (1.1 ± 0.1 found here agrees with 1.1 reported by Dessy et al.30). It is interesting that cyclic voltammetry shows that the cathodic peak encompasses two electron reductions to a dianion, but coulometry shows an integrated total of one electron per molecule. This is strong support for the homoassociation of each [2Fe-2S]2− dianion with another [2Fe-2S] neutral molecule to form a [4Fe-4S]2− dianion, with an average of one electron per [2Fe-2S] molecule in bulk electrolysis.

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Figure 8. IR spectroelectrochemistry of (μ-SMe)2Fe2(CO)6 in an Ar saturated solution of 0.3 M TBAPF6 in CH3CN. Top: dark blue curve is carbonyl absorptions following complete reduction of the starting material at ∼0.2 V negative of the cathodic CV peak. The lighter colored lines are repeated IR every 5 min for 20 min showing slow return of the starting material absorptions. After 20 min, the material was reoxidized at ∼0.2 V, and the IR showed 97−100% recovery of the starting material. Bottom: computational gas phase carbonyl absorptions for the dimer shown in Scheme 3.

DISCUSSION AND CONCLUSIONS The [2Fe-2S] cluster core in these organometallic compounds is quite malleable. Both oxidation and reduction induce substantial geometric rearrangements to optimize the bonding and stabilize the cluster charge. In the cation, the primary geometric rearrangement is the shift of a carbonyl ligand to a semibridging position. This structure is evidenced by the onset of gas-phase ionization energies in combination with computations. The onset of oxidation in solution is also consistent with this structure, but unfortunately, the oxidation is fully irreversible and does not yield additional information. The irreversibility may be due to the rotation of a carbonyl to a semibridging site that opens up a coordination site susceptible to nucleophilic attack. Immediate possibilities include the solvent, the counteranion in solution, another cluster molecule, or surface reactions. Reduction opens up the [2Fe-2S] butterfly. The first electron reduction elongates the Fe−Fe distance toward the square [2Fe-2S] arrangement as typically observed for metal carbonyl complexes. However, the second electron reduction strongly favors a structure in which one Fe−S bond is broken, and a terminal carbonyl moves to a bridging position. This structure is sufficiently effective at stabilizing the negative charge that the potential for the second electron addition is less negative than the potential for the first electron addition (potential inversion). The extensive geometric rearrangement from the anion to the dianion leads to a comparatively slow electron transfer rate, as evidenced by the CV scan rate studies. This geometric stabilization of the dianion appears to be characteristic of other hydrogenase mimics, with the difference in behavior of the first cathodic peak largely dependent on the rate of geometric rearrangement and second electron transfer. For the well-known (μ-1,3-propanedithiolato)Fe2(CO)6 compound (PDT), Capon et al. found this bridging carbonyl structure with one broken Fe−S bond to be most stable for the dianion,52 but the first cathodic peak is primarily a one-electron reduction due to slow rearrangement. The (μ-1,2ethanedithiolato)Fe2(CO)6 compound (EDT) has an initial two-electron reduction at a 0.02 V/s scan rate, but at the faster 20 V/s rate, the reorganization to the stable dianion is outpaced, and the CV peak becomes a one-electron

report in the region from 1850−2100 cm−1 (the previous report did not scan below 1850 cm−1).27 The spectrum closely matches that obtained from reduction of (μ-SEt)2Fe2(CO)6 through the full region.31 Nearly complete reduction of the starting compound is observed as shown by the loss of the strong absorption of the neutral compound at 2035 cm−1. Over a period of 20 min, about 20% of the absorption intensity of the neutral compound reappeared, presumably due to slow diffusion in the IR spectroelectrochemistry cell. Shifting the potential to −1.3 V, which is positive of the first reduction peak, did not produce any current, indicating the absence of [2Fe-2S] anions. The IR did not change at this potential. Most significantly, after 20 min the solution was reoxidized at −0.2 V, which is positive of the return anodic peaks, and the IR showed 97−100% recovery of the starting material. The experiment was cycled back and forth from fully reduced to fully reoxidized species multiple times. All of the absorptions produced by electrolysis rise and fall together indicating a single molecular species. Also shown are the calculated carbonyl absorptions for the homoassociated [4Fe-4S]2− dianion pictured in Scheme 3. The experimental IR and the calculations have several parallels. Most important, the experiment shows a broad absorption in the bridging carbonyl region, and the optimum computational structure of the homoassociated [4Fe-4S]2− species has bridging carbonyl ligands. Both the experiment and the calculations also have a cluster of absorptions from about 1850−1950 cm−1, a prominent single absorption over 1950 cm−1, and weaker absorption around 2025 cm−1. The number of absorptions is more representative of a [4Fe-4S] species than a single [2Fe2S] species. The IR spectrum is very similar to that reported for the dianion of the (μ-biphenyl-2,2′-dithiolato)Fe 2(CO)6 5016

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reduction.47 Reduction of the MeMe compound in this study remains a two-electron process up to 50 V/s scan rate (Figure S2, Supporting Information). The (μ-1,2-benzenedithiolato)Fe2(CO)6 compound (BDT) is different in that the first electron reduction causes distortion directly toward the stabilized dianion structure, and consequently, the second electron reduction is relatively fast. The dianion is susceptible to other forms of stabilization. The present study shows that the MeMe compound, in the absence of other species to better stabilize the dianion charge, will homoassociate the dianion with another MeMe molecule to form a [4Fe-4S]2− dianion. The additional stabilization is evidenced by the return oxidations that are at potentials substantially more positive than those of the initial reduction. Interestingly, these return oxidations regenerate the starting compound so that the total reduction/reoxidation cycle shows chemical reversibility. The absence of irreversible processes such as carbonyl ligand loss, disproportionation, or other chemical pathways with reduction is significant since this MeMe compound would seem as likely to be susceptible to these other processes as many other hydrogenase active-site mimics. Evidence that homoassociation is responsible for stabilization of the dianion comes from the scan rate studies, chronocoulometry, and IR spectroelectrochemistry, all of which are consistent with the computations. Coulometry is particulary interesting in this case because the CV shows that each molecule undergoes two-electron reductions in the first cathodic peak, but because of homoassociation, the net result of bulk electrolysis is one electron per molecule. Thus, earlier works that differ in reporting a one-electron30 or a twoelectron29,32 reduction are both correct depending on the method of determination. Differences in chemical reversibility also depend on the mechanisms available to stabilize the dianion charge. Homoassociation has been identified for other complexes.57 Both the PDT31 and EDT54,56 compounds are known to form [4Fe-4S]2− dianions with different structures. The BDT compound does not show a tendency for homoassociation, and as a consequence, a chemically reversible two-electron reduction is observed with little separation between the cathodic and anodic peaks at normal scan rates. The crystal structure of the BDT [2Fe-2S]2− dianion confirms the structure with the bridging carbonyl and elongated Fe−S distance.53 The broader impact and insight provided by this work is that the reductive opening of the [2Fe-2S] core exhibited by these compounds is likely to be a general feature of this class of complexes, modified primarily by the rate of the geometric transformation. The PDT compound is unique among the compounds discussed here in having the slowest transformation to the stable dianion structure. This suggests a structural benefit to the three-atom linker between the sulfur atoms in the active sites of [FeFe] hydrogenases because stabilization of the dianion by this structural distortion diminishes the chemical potential for protonation leading to reductive catalysis. The second point is that associative processes with reduction may dominate over dissociative processes. Homoassociation further stabilizes the dianion and diminishes the chemical potential for protonation leading to reductive catalysis. Homoassociation is not possible in the enzyme, but the [2Fe-2S] active site in the enzyme is instead already associated with a [4Fe-4S] cubane cluster through a Fe−S(cysteine)−Fe linkage to form the Hcluster, similar to the SMe linkage depicted in Scheme 3 for the MeMe [4Fe-4S]2− species. As pointed out previously for the H-

cluster58 and found here for the MeMe compound, cluster associaton with the [2Fe-2S] active site is electronically linked and modulates the reduction chemistry.

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EXPERIMENTAL SECTION

General Procedures. All reactions were performed under an Ar atmosphere using standard Schlenk line techniques. All glassware was dried in a 110 °C oven overnight before use. Commercially available chemicals iron pentacarbonyl (Aldrich, >99.99%), 2-iodomethane (Aldrich, 99%), and DrySolv THF were used as received. Solvents were deoxygenated by sparging with Ar for 30 min prior to use or used directly from a MBraun solvent purification system (EMD Omnisolv). All reactions were monitored by solution IR spectroscopy of the CO stretching region and/or by TLC on Merck aluminum TLC sheets. Work up and chromatographic separations were carried out in air using silica gel columns with pentane or hexanes as eluents. Product yields were calculated based on utilized (μ-S)2Fe2(CO)6. Infrared spectra were collected in mineral oil or in hexane on round sodium chloride cells using a Nicolet 380 FT-IR spectrophotometer. 1H and 13 C NMR spectra were acquired at room temperature on a Bruker AVIII 400 spectrometer using CDCl3 as solvent. The starting compound, (μ-S2)Fe2(CO)6, was prepared following literature methods with some modification.11 Preparation of (μ-SMe)2Fe2(CO)6. The (μ-SMe)2Fe2(CO)6 complex was prepared following the literature method.11 (μ-S2)Fe2(CO)6 (1.02 g (3 mmol)) was added to a 250 mL round-bottomed flask, purged with Ar for 30 min, and 60 mL of dry THF added to the flask via cannula. The mixture was stirred and cooled to −78 °C in a dry ice/acetone bath. Lithium triethylborohydride in THF (1 M, Aldrich, 99.8%) was slowly added to the reaction flask by adding 0.5 mL of lithium triethylborohydride every 5 min with stirring for a total of 6 mL (6 mmol). The reaction mixture was stirred at −78 °C for 30 min, and then, 0.9 mL (6 mmol) of iodomethane was added via syringe. The reaction was stirred at −78 °C for 20 min, then warmed to room temperature and stirred for 3 h, and the solvent removed by rotary evaporation. The dark red crude product was purified by a silica gel chromatographic column, 2.5 × 25 cm2, using hexane as eluent. The (aMe)(eMe) isomer was isolated as a major product. Yield 0.35 g (1 mmol) 33%; Rf (hexane) 0.52; IR (hexane, CO stretching, cm−1): 2071 (w), 2035 (vs), 1999 (m), and 1989 (m). 1H NMR (CDCl3) δ (ppm), 1.24 (s, CH3 a), 1.61 (s, CH3 e). Electrochemistry. Sources and treatment of solvent and electrolyte are the same as those reported earlier.59 Electrochemical procedures including the determination and compensation of solution resistance have also been reported.59 Cyclic voltammetry experiments were carried out using a Gamry Reference 3000 potentiostat/ galvanostat. A standard 3-electrode system was utilized including a Ag/0.01 M AgNO3 reference electrode, a glassy-carbon working electrode (GCE) with 3 mm diameter determined to have an area of 0.071 cm2, and a Pt wire auxiliary electrode. The reference electrode was corrected against a 1.0 mM solution of ferrocene in acetonitrile before and after each experiment, and all potentials are reported against Fc+/Fc.60 General conditions for the cyclic voltammetry experiments included approximately 1 mM of the diiron compound in acetonitrile containing 0.1 M n-Bu4NPF6 (TBAPF6) as supporting electrolyte, ambient temperature, and a minimum of 30 s between scans with stirring under N2 or Ar. The electrode was polished with 0.05 μm alumina in deionized water on a felt surface prior to each experiment and sonicated in methanol followed by acetonitrile before introduction to the electrochemical cell. Simulations were carried out with the program DigiElch.61 All cyclic voltammgrams are background corrected. IR spectroelectrochemistry experiments were conducted in a cell of a design similar to that in the literature.62 The IR spectra were collected on a Thermo Nicolet Avatar ESP 380 FT-IR spectromenter. Controlled potential coulometry experiments were carried out in the same cell, and the IR spectra were used to confirm complete reduction and complete reoxidation of the sample. The number of coulombs passed for a known concentration of solution were calibrated against 5017

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(2) Lubitz, W.; Ogata, H.; Rudiger, O.; Reijerse, E. Chem. Rev. 2014, 114, 4081−4148. (3) Nicolet, Y.; Piras, C.; Legrand, P.; Hatchikian, C.; FontecillaCamps, J. C. Struct. Fold Des. 1999, 7, 13−23. (4) Fontecilla-Camps, J. C.; Volbeda, A.; Cavazza, C.; Nicolet, Y. Chem. Rev. 2007, 107, 4273−4303. (5) Armstrong, F. A. Electroanal. Chem. 2014, 25, 33−103. (6) Wang, N.; Wang, M.; Chen, L.; Sun, L. Dalton Trans. 2013, 42, 12059−12071. (7) Wright, J. A.; Pickett, C. J. ChemCatChem 2012, 4, 1723−1724. (8) Wei, C. H.; Dahl, L. F. Inorg. Chem. 1965, 4, 1−11. (9) Farrugia, L. J.; Evans, C.; Senn, H. M.; Hanninen, M. M.; Sillanpaa, R. Organometallics 2012, 31, 2559−2570. (10) Eremenko, I. L.; Berke, H.; Van, d. Z.; Kolobkov, B. I.; Novotortsev, V. M. J. Organomet. Chem. 1994, 471, 123−132. (11) Seyferth, D.; Henderson, R. S.; Song, L.-C. Organometallics 1982, 1, 125−133. (12) Seyferth, D.; Henderson, R. S.; Song, L-C.; Womack, G. B. J. Organomet. Chem. 1985, 292, 9−17. (13) Seyferth, D.; Henderson, R. S.; Gallagher, M. K. J. Organomet. Chem. 1980, 193, C75−C78. (14) Seyferth, D.; Song, L.-C.; Henderson, R. S. J. Am. Chem. Soc. 1981, 103, 5103−5107. (15) Gloaguen, F.; Rauchfuss, T. B. Chem. Soc. Rev. 2009, 38, 100− 108. (16) 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. (17) Schilter, D.; Rauchfuss, T. B. Angew. Chem., Int. Ed. 2013, 52, 13518−13520. (18) Capon, J.; Ezzaher, S.; Gloaguen, F.; Petillon, F. Y.; Schollhammer, P.; Talarmin, J. Chem.Eur. J. 2008, 14, 1954−1964. (19) Singleton, M. L.; Bhuvanesh, N.; Reibenspies, J. H.; Darensbourg, M. Y. Angew. Chem., Int. Ed. 2008, 47, 9492−9495. (20) Hsieh, C.; Erdem, O. F.; Harman, S. D.; Singleton, M. L.; Reijerse, E.; Lubitz, W.; Popescu, C. V.; Reibenspies, J. H.; Brothers, S. M.; Hall, M. B.; Darensbourg, M. Y. J. Am. Chem. Soc. 2012, 134, 13089−13102. (21) Eisch, J. J.; King, R. B. Organometallic Syntheses (Transition-Metal Compounds, Vol. 1); Academic Press: New York, 1965. (22) King, R. B. J. Am. Chem. Soc. 1962, 84, 2460. (23) Crow, J. P.; Cullen, W. R. Can. J. Chem. 1971, 49, 2948−2952. (24) De Beer, J. A.; Haines, R. J.; Greatrex, R.; Greenwood, N. N. J. Chem. Soc. A 1971, 3271−3282. (25) Adams, R. D.; Cotton, F. A.; Cullen, W. R.; Hunter, D. L.; Mihichuk, L. Inorg. Chem. 1975, 14, 1395−1399. (26) Hall, M. B.; Fenske, R. F.; Dahl, L. F. Inorg. Chem. 1975, 14, 3103−3117. (27) Dessy, R. E.; Wieczorek, L. J. Am. Chem. Soc. 1969, 91, 4963− 4974. (28) Mathieu, R.; Poilblanc, R. J. Organomet. Chem. 1977, 142, 351− 355. (29) Mathieu, R.; Poilblanc, R.; Lemoine, P.; Gross, M. J. Organomet. Chem. 1979, 165, 243−252. (30) Dessy, R. E.; Stary, F. E.; King, R. B.; Waldrop, M. J. Am. Chem. Soc. 1966, 88, 471−476. (31) de Carcer, I. A.; DiPasquale, A.; Rheingold, A. L.; Heinekey, D. M. Inorg. Chem. 2006, 45, 8000−8002. (32) Darchen, A.; Mousser, H.; Patin, H. Chem. Commun. 1988, 968−970. (33) Borg, S. J.; Ibrahim, S. K.; Pickett, C. J.; Best, S. P. C. R. Chim. 2008, 11, 852−860. (34) Ortega-Alfaro, M. C.; Hernández, N.; Cerna, I.; López-Cortés, J. G.; Gómez, E.; Toscano, R. A.; Alvarez-Toledano, C. J. Organomet. Chem. 2004, 689, 885−893. (35) Lyon, E. J.; Georgakaki, I. P.; Reibenspies, J. H.; Darensbourg, M. Y. Angew. Chem., Int. Ed. 1999, 38, 3178−3180. (36) Allen, F. H. Acta Crystallogr., Sect. B 2002, 58, 380−388.

the number of coulombs passed for a known concentration of ferrocene. The coulomb traces vs time were similar for oxidation and reduction. Gas-Phase UV Photoelectron Spectroscopy. Photoelectron spectra were recorded using an instrument that features a 360 mm radius hemispherical analyzer63 with custom-designed photon source, sample cells, detection, and control electronics as described previously.64 Data collection and processing were carried out as described previously.50 The MeMe compound sublimed at room temperature. Density Functional Theory Computations. DFT calculations were performed using Amsterdam Density Functional (ADF) software, version 2013.01.65−67 Initial geometry optimizations and frequency calculations were carried out in the gas phase using the local density approximation (LDA)68 and the Vosko−Wilk−Nusair (VWN) functional69 with the Stoll correction implemented.70 The basis set was triple-ζ with one polarization function (TZP). Relativistic corrections were included according to the zeroth-order relativistic approximation (ZORA),65 available in the ADF package. Subsequently, the geometries and energies were refined in solution with the PBE functional71 with dispersion corrections according to the method of Grimme using the BJ damping function (PBE-D3-BJ).72 Solvation Gibbs energies are estimated by the conductor-like screening model73 (COSMO) of solvation using default parameters for acetonitrile. Calculations of the standard oxidation and reduction potentials and Gibbs energies and equilibrium constants for reactions included the optimized electronic energies in solvent, the zero-point vibrational energies unscaled from harmonic frequency calculations, and thermal enthalpy and entropy contributions in solution at 298.15 K. The solution translational and rotational entropy was estimated as described before50 but regardless of the definition has little effect on the calculated potentials. Figures of the optimized geometries and molecular orbitals were created with program Visual Molecular Dynamics.74

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ASSOCIATED CONTENT

S Supporting Information *

Calculated relative energies of isomers, valence orbitals, and energies, details of the CV simulations, and DFT optimized geometries for convenient visualization. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Authors

*(K.J.H.) E-mail: [email protected]. *(D.L.L.) E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The support of the National Science Foundation (CHE0527003 and CHE-1111570) and the Commission on Higher Education, Thailand (to O.I. and K.J.H., CHE-PHD-SW-2549) are gratefully acknowledged. O.I. acknowledges Professor John H. Enemark for catalyzing her international adventures. Additionally, special thanks to Professor Clifford Kubiak and members of his research group for aid in the construction of the IR-SEC cell.

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ABBREVIATIONS ADF, Amsterdam density functional; CV, cyclic voltammetry; DFT, density functional theory; Me, methyl; TBAPF6, nBu4NPF6; IR-SEC, infrared spectroelectrochemistry

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REFERENCES

(1) Tard, C.; Pickett, C. J. Chem. Rev. 2009, 109, 2245−2274. 5018

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dx.doi.org/10.1021/om5004122 | Organometallics 2014, 33, 5009−5019