FeMo Heterobimetallic Dithiolate Complexes: Investigation of Their

Article Cite This: Inorg. Chem. XXXX, XXX, XXX−XXX

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FeMo Heterobimetallic Dithiolate Complexes: Investigation of Their Electron Transfer Chemistry and Reactivity toward Acids, a Density Functional Theory Rationalization Solène Bouchard,† Maurizio Bruschi,*,‡ Luca De Gioia,§ Christine Le Roy,† François Y. Pétillon,† Philippe Schollhammer,*,† and Jean Talarmin*,†

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†

UMR CNRS 6521 “Chimie, Electrochimie Moléculaires et Chimie Analytique”, Université de Bretagne Occidentale, UFR Sciences et Techniques, CS 93837, 29238 Brest-Cedex 3, France ‡ Department of Earth and Environmental Sciences, University of Milano-Bicocca, Piazza della Scienza 1, 20126 Milan, Italy § Department of Biotechnology and Bioscience, University of Milano-Bicocca, Piazza della Scienza 2, 20126 Milan, Italy S Supporting Information *

ABSTRACT: The electrochemical behavior of complexes [FeMo(CO)5(κ2-dppe)(μ-pdt)] (1) and [FeMo(CO)4(MeCN)(κ2-dppe)(μpdt)] (2), in the absence and in the presence of acid, has been investigated. The reduction of 1 follows at slow scan rates, in CH2Cl2− [NBu4][PF6] and acid-free media, an ECrevE mechanism that is supported by cyclic voltammetry (CV) experiments and digital CV simulations. In MeCN−[NBu4][PF6], the electrochemical reduction of 1 is the same as in dichloromethane and follows an ECE mechanism at slow scan rates, but with a positive shift of the redox potentials. In contrast, the oxidation of 1 is strongly solvent-dependent. In dichloromethane, the oxidation of 1 is reversible and involves a single electron, while in acetonitrile, it is irreversible at moderate and slow scan rates (v ≤ ca. 1 V s−1), and some chemical reversibility is apparent at higher scan rates (v = 10 V s−1). Density functional theory calculations revealed that the chemical step in the ECrevE mechanism corresponds to the dissociation of one PPh2 end of the diphosphine ligand and the transfer of the semibridging CO to the Fe atom, similarly to the mechanism observed in the FeFe analogue complex. However, in the case of 1, the subsequent coordination of the phosphine ligand to the other metal is an unfavorable process.

1. INTRODUCTION Nature widely uses metal−sulfur assemblies for the transfer of electrons, and for the reduction (N2, H+) or the oxidation (H2) of various substrates.1−9 Whereas dihydrogen production and uptake are most efficiently catalyzed by the homometallic [FeFe] hydrogenases,4−6 the most competent enzyme for the energy costly reduction of dinitrogen to ammonia is the molybdenum nitrogenase,7 which possesses a single molybdenum atom at the active site, the (R-homocitrate){MoFe7S9C} cluster (the FeMo cofactor or FeMo-co).10−13 Although recent biochemical experiments14−16 and density functional theory (DFT) calculations17 tend to favor iron atoms as the binding site of the substrate N2 (and of the inhibitor/substrate CO),18 our understanding of where and how N2 is reduced by the Mo nitrogenase is still limited and the possible involvement of molybdenum in the processes is not ruled out.19−23 Model studies involving either Fe or Mo complexes have extended our knowledge on how the NN triple bond, or the NNR double bond, is progressively reduced at the metal center(s) of synthetic complexes,22−33 but only rare examples of heterometallic systems are known. At any rate, because the way the molybdenum atom improves the efficiency of N2 reduction at the expense of H2 production (when compared to the all-iron nitrogenase) is unknown,22,23 it is interesting to © XXXX American Chemical Society

compare the reactivity of analogue complexes of Mo and Fe in order to try and understand the differences induced by changing the metal centers.34,35 We have recently described the synthesis and characterization of the novel FeMo complex [FeMo(CO)5(κ2-dppe)(μpdt)] (1) (pdt = propanedithiolate)36 which is an heteronuclear analogue of the diiron compound [Fe2(CO)4(κ2dppe)(μ-pdt)] reported earlier.37 In particular, Mössbauer and DFT analysis of the electronic structure revealed the unsymmetrical charge distribution of 1 in which the FeIIMo0 resonant form gives a significant contribution to proper representation of its electronic structure. We have now investigated the electrochemical behavior of complex 1 and of the MeCN-substituted derivative [FeMo(CO)4(MeCN)(κ2dppe)(μ-pdt)] (2) in the absence and in the presence of acid. The electrochemical results have been complemented by DFT calculations, which have been useful to unveil the mechanistic details of the electrochemical processes. Received: October 11, 2018

A

DOI: 10.1021/acs.inorgchem.8b02861 Inorg. Chem. XXXX, XXX, XXX−XXX

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Inorganic Chemistry

2. RESULTS AND DISCUSSION 2.1. Electrochemistry of [FeMo(CO)5(κ2-dppe)(μ-pdt)] (1) in Acid-Free Media. 2.1.1. In CH2Cl2−[NBu4][PF6]. The cyclic voltammetry (CV) of [FeMo(CO)5(κ2-dppe)(μ-pdt)] (1)36 in CH2Cl2−[NBu4][PF6] shows that the complex undergoes reversible or partially reversible redox processes (Figure 1, Table 1). The cyclic voltammograms under an

Figure 2. Scan rate dependence of the current functions of the reduction and of the oxidation of [FeMo(CO)5(κ2-dppe)(μ-pdt)] (1, 1.03 mM) in CH2Cl2−[NBu4][PF6].

indicates, by comparison with that of the oxidation, that more than one electron is transferred on the longest time scale. Under these conditions, a reduction product is detected by its oxidation at −1.43 V on the reverse scan (Figure 1). The CV in Figure 3 shows that the oxidation process at −1.43 V leads to the regeneration of some [FeMo(CO)5(κ2-

Figure 1. CV of [FeMo(CO)5(κ2-dppe)(μ-pdt)] (1, 1.41 mM) in CH2Cl2−[NBu4][PF6] under Ar (vitreous carbon electrode, v = 0.2 V s−1, potential in V/Fc+/Fc).

Table 1. Redox Potentials [V vs E1/2 (Fc+/Fc)] of the [FeMo(CO)4(L)(κ2-dppe)(μ-pdt)] Complexes (L = CO: 1; L = MeCN: 2) complex 2

[FeMo(CO)5(κ -dppe)(μ-pdt)] (1) [FeMo(CO)4(MeCN)(κ2 -dppe)(μ-pdt)] (2)

solvent

Ered 1/2 (V)

Eox 1/2 (product)a (V)

CH2Cl2

−1.84

−1.43

MeCN MeCN

−1.70 −2.03d

−1.38b −1.86c

b

Eox 1/2 (V) 0.12 0.18c −0.13

a

Oxidation peak detected on the reverse scan of the CV of 1 or 2 (see text). bIrreversible. cReversible, v = 10 V s−1. dIrreversible, v = 10 V s−1.

Figure 3. CV of [FeMo(CO)5(κ2-dppe)(μ-pdt)] (1, 1.34 mM) in CH2Cl2−[NBu4][PF6] showing the second negative-going scans in different potential ranges with scan reversal at (black curve): −1.08 V and (red curve): −1.58 V (vitreous carbon electrode, v = 0.2 V s−1, potential in V/Fc+/Fc).

atmosphere of dinitrogen are exactly the same as those obtained under Ar (Figure 1), which demonstrates that neither the neutral complex 1 nor its oxidized or reduced forms react with N2 on the CV timescale under the present experimental conditions. 38 The oxidation of 1 at Eox is reversible and 1/2 = 0.12 V 39,40 involves a single electron, as evidenced by the comparison 1/2 1/2 of its current function (iox p /v CD0 ; C is the concentration of the complex and D0 is its diffusion coefficient) with that of the diiron compound [Fe2(CO)4(IMe−CH2−IMe)(μ-pdt)] (IMe = 1-methylimidazol-2-ylidene) shown previously to undergo a reversible one-electron oxidation41 (Figure S1, Supporting Information). The reversible oxidation of 1 may thus be taken as a reference for a one-electron transfer for this series of compounds. Other irreversible oxidation processes observed at more positive potentials are not shown in Figure 1 and were not investigated. The electrochemical reduction of 1 looks more interesting than the simple one-electron oxidation. At a slow scan rate (v = 0.2 V s−1, Figure 1), the first reduction at Ered 1/2 = −1.84 V is a partially reversible step whose current function (Figure 2)

dppe)(μ-pdt)] (1). Indeed, the amount of starting complex 1 present at the electrode surface on the second negative-going scan is larger when the oxidation peak at −1.43 V is traversed before scan reversal (Figure 3, black line) than when it is not traversed (Figure 3, red line). Keeping in mind the fact that the time required to scan back and forth from −1.58 to −1.08 V (i.e., 5 s at v = 0.2 V s−1) allows for diffusion of complex 1 from the bulk to the electrode surface, Figure 3 suggests that the irreversible oxidation at −1.43 V regenerates the starting complex 1, at least partly. The pattern consisting of the system at Ered 1/2 = −1.84 V, of the small reduction peak around −2.2 V, and of the large oxidation peak at −1.43 V on the return scan, resembles what is observed when the two steps of an EE process (with inverted potentials)42−49 are kinetically discriminated.48,49 Because it is known that several hexacarbonyl diiron dithiolate complexes related to 1 undergo such two-electron reduction processes,50−66 we examined whether this is also true for 1. In B


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Inorganic Chemistry

oxidation. Increasing the scan rate does not allow the detection of a return peak for the P2−/P− system: indeed, the P2− oxidation is suppressed before its chemical reversibility can be observed. This indicates that the P− → 1− reaction is faster than the opposite one. Digital CV simulations with DigiElch67 were conducted in order to assess whether the reactions in Scheme 1 are consistent with the experimental results. That it is the case is confirmed by the fair agreement (Figure 4) between the experimental data and the curves obtained by digital simulation of the mechanism in Scheme 1 using the parameters shown in Table 2. It must be noted that the small

the case of a kinetically discriminated EE process, the small peak at −2.2 V would arise from the reduction of the anion 1− to the dianion 12−, while the peak at −1.43 V would correspond to the two-electron oxidation of 12− back to 1. However, the scan rate dependence of the peak current ratio red1 is not consistent with such an EE mechanism.38 ired2 p /ip Indeed, the fact that the first reduction becomes chemically more reversible at fast scan rates (Figure S2, Supporting Information) signifies that the amount of 1− present at the electrode surface increases as v increases. If the reduction at −2.2 V was due to the 1 − → 1 2− conversion, the red1 corresponding peak (and ired2 p /ip ) should increase with v. The opposite is observed. At the same time, the oxidation peak at −1.43 V is suppressed upon increasing v, even when the cathodic scan is extended beyond the peak at −2.2 V (Figures S2 and S3, Supporting Information). All these features are consistent with the reduction of 1 following an ECrevE mechanism in which the formation of the product P− (Scheme 1) is prevented at fast scan rates.

Table 2. Parameters Used for the Simulations of the Reduction of 1 According to an ECrevE Mechanism (the Figures in Italics Are Calculated by the DigiElch Software)a redox steps

E1/2 (V)

α

ks (cm/s)

1 + 1e− = 1− P− + 1e− = P2− chemical reactions

−1.84 −1.43 Keq

0.5 0.5 kf

0.05 0.005 kb

10 100

166.7 1.74 × 10−5

1− = P− 1− + P− = 1 + P2−

Scheme 1. Proposed Mechanism for the Reduction of [FeMo(CO)5(κ2-dppe)(μ-pdt)] in CH2Cl2−[NBu4][PF6] Including an ECrevE Process (Black Arrows) and a Homogeneous Redox Reaction between the Anionic Species (P− + 1− = P2− + 1) (Red Arrows)

0.06 5.76 × 106

D0 = 4.9 × 10−6 cm2/s; C0 = 1.34 × 10−3 mol/L; A = 0.072 cm2; Ru = 100 Ω; Cd = 1.5 × 10−6 F. a

reduction at −2.2 V, which probably arises from a secondary (decomposition) product, was not included in the simulations. It is also interesting to note that, in order to obtain a satisfying agreement between the experimental and simulated curves, the rate constant of the heterogeneous electron transfer for the system at −1.43 V (P2−/P−) has to be 10 times smaller than that of the 1/1− couple (Table 2). This suggests that a chemical reaction may be concurrent with the transfer of the second electron. The reduction of the FeMo complex 1 is thus quite different from that of the diiron analogue, [Fe2(CO)4(κ2-dppe)(μpdt)], which gives rise to an electron transfer chain (ETC) catalytic isomerization to the (μ-dppe) compound.68 In the mechanism proposed for this isomerism, one CO ligand moves from one Fe center to the adjacent one, which in turn causes the partial decoordination of the dppe chelate of the anionic species. The structural modification is completed by the coordination of the dppe in a bridging mode, and by an electron transfer producing the neutral μ-dppe isomer. Theoretical calculations were undertaken in order to look into the possible structural consequences of the reduction of 1. The comparison of the structures of the neutral and anionic

That the oxidation of P2− at −1.43 V regenerates the starting complex results from the reversibility of the intervening chemical step. Thus, the oxidation of P2− produces 1− (via P−), which is oxidized to 1 at a potential more negative than the potential of the P2− oxidation. In addition, because 1− oxidizes at Eox 1/2 = −1.84 V, that is, a potential 0.4 V more negative than the reduction of P − (ca. −1.43 V), a homogeneous redox reaction between the anionic species (Scheme 1, red arrows) may contribute to the overall P2− ↔ 1 interconversion process. Therefore, the oxidation of P2− appears as a chemically irreversible, overall two-electron

Figure 4. Experimental and simulated cyclic voltammograms of [FeMo(CO)5(κ2-dppe)(μ-pdt)] (1, 1.34 mM) in CH2Cl2−[NBu4][PF6] in different potential ranges (vitreous carbon electrode, v = 0.2 V s−1, potential in V/Fc+/Fc). C


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Inorganic Chemistry complexes optimized by DFT in CH2Cl2 implicit solvent shows that reduction results in an elongation of the Fe−Mo distance from 2.853 to 3.280 Å, and of the distance between the iron atom and the carbon atom of the semibridging CO from 2.549 Å in the neutral complex to 3.112 Å in the anion (Figure S4, Supporting Information), as expected from the antibonding character of the lowest unoccupied molecular orbital (LUMO) of 1 (Figure 5). The Fe−S and Mo−S

Figure 6. Schematic representation of the structures of the proposed P− (left) and P2− (right) species with selected distances (in Å) calculated at the BP86/def-TZVP level of theory in CH2Cl2.

Table 3. Relative Gibbs Free Energies (in kcal mol−1) of the Different Isomers of 1, 1−, and 12− in vacuum and in CH2Cl2 and CH3CN Solvents 1 1(κ1-dppe)P 1(μ-dppe) 1− 1−(κ1-dppe)P− 1−(μ-dppe) 12− 12−(κ1-dppe)P2− 12−(μ-dppe)

Figure 5. View of the LUMO of 1 (optimized structure from DFT calculations; H: white, C: gray, O: red, P: light green, Mo: dark green, S: yellow, Fe: yellow).

distances are also affected by the reduction, going from 2.306 Å (Fe−S) and 2.565 Å (Mo−S) in the neutral complex to 2.329 Å (Fe−S) and 2.636 Å (Mo−S) in the anion. The reduction of 1− to the corresponding dianion (12−) results in a further elongation of the Fe−Mo and semibridging OC−Fe distances up to 3.515 and 3.265 Å, respectively, as well as to the elongation of the Fe−S and Mo−S distances up to 2.373 and 2.651 Å, respectively (Figure S4). The reduction of 1 also affects the relative stability of the apical−basal and basal−basal isomers. In fact, in the neutral complex 1, the basal−basal isomer is significantly more stable than the apical− basal one (Table S1, Supporting Information), whereas the Gibbs free energy difference is reduced to about 2 and 1 kcal mol−1 in 1− and 12−, respectively. The redox potential calculated in CH2Cl2 for the 1 → 1− process is equal to −1.90 V, which is in very good agreement with the experimental value. On the other hand, the redox potential calculated for the second reduction, from 1− to 12− is as low as −3.11 V. Even if this value is expected to be overestimated, because of the large negative charge of the reactant, it is too low to fit with the experimental observation, that is, the second reduction potential is more positive than the first one. Therefore, the geometrical rearrangement occurring from 1 to 1− and 12− should not correspond to the chemical step in the ECrevE mechanism. Consequently, an extensive search on the potential energy surface (PES) of 1− was carried out in order to identify other energy minima that can explain such a reversible chemical step. Indeed, a stable isomer characterized by the dissociation of one of the two PPh2 ends of the chelating diphosphine ligand, and the migration of the semibridging CO to the Fe atom, was located on the PES. Such isomer, which should correspond to the species P− (Figure 6), is less stable than 1− by 3 kcal mol−1 in vacuum, but it becomes slightly more stable than 1− when geometry optimization is carried out in CH2Cl2 and CH3CN implicit solvents (Table 3). Notably, the dianionic species formed after the second reduction, which is assigned to P2− (Figure 6), is more stable

vacuum

CH2Cl2

CH3CN

0.0 23.0 3.8 0.0 3.0 6.6 19.5 0.0 25.6

0.0 22.6 4.4 3.0 0.0 8.4 31.1 0.0 30.9

0.0 22.6 4.4 4.2 0.0 9.0 33.2 0.0 31.5

than that obtained by the reduction of 1− by as much as 20, 31, and 33 kcal mol−1 in vacuum, CH2Cl2 and CH3CN solvents, respectively. The stabilization of P2− with respect to 12− can be explained by considering the accumulation of charge on the Fe atom in the dianion. Indeed, population analysis shows that the bielectronic reduction from 1 to 12− corresponds to a change in the formal redox state from Fe(II)Mo(0), as previously discussed,36 to Fe(0)Mo(0), where the Fe atom is stabilized by the replacement of a strong σ-donor phosphine ligand with a strong π-accepting CO ligand. The redox potential calculated for the reduction from P− to P2− in CH2Cl2 is equal to −1.89 V, a value still lower than the experimental one (−1.4 V) by about 0.5 V, but compatible with the bielectronic reduction of 1. It is interesting to note that the chemical process described above corresponds to the first step proposed for the ETC catalytic isomerization of [Fe2(CO)4(κ2-dppe)(μ-pdt)] to the (μ-dppe) isomer.68 The second chemical step in such a process is the binding of the dangling phosphorus atom to the coordinatively unsaturated iron center to give the (μ-dppe) compound, which is much more stable than the parent species. In the reduction of 1, this second chemical step does not occur because the (μ-dppe) isomer 1(μ-dppe) in Table 3, which has been identified as a genuine minimum on the PES in our calculations, is less stable than 1 (Table 3). Other minima on the PES have been identified, being characterized by the dissociation of either one S−Fe or S−Mo bonds, which have been proposed to be involved in the reduction of [Fe2(CO)6{μ-SCH2XCH2S}] (X = CH2, NR) complexes. However, all these isomers are significantly less stable than 1 or its reduced species and therefore should not be involved in the reduction of the parent compound. 2.1.2. In MeCN−[NBu4][PF6]. The electrochemical reduction of 1 in acetonitrile is the same as in dichloromethane, but with D


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mechanism corresponds to the substitution of one basal CO (or of both basal COs) of the cation 1+ by acetonitrile molecule(s). The replacement of carbonyl(s) by better electron donor ligand(s) would favor the transfer of a second electron, as observed. In fact, theoretical calculations support this hypothesis. The relative stability of the CH3CN adduct obtained by substituting the basal/apical (Mo)−CO with an acetonitrile molecule can be evaluated by considering the Gibbs free energy (ΔGreact) of the reaction 1x + CH3CN → 2x + CO (x = 0, +1, +2). Interestingly, as shown in Table S2 of Supporting Information, the stability of the CH3CN adduct significantly increases with the positive charge of the complex. In addition, the relative stability of the isomer with the basal CH3CN ligand coordinated to Mo also increases with respect to that in which CH3CN is apical by augmenting the positive charge of the complex (Table S2). In particular, the isomer with apical CH3CN is more stable than the basal one by about 7 and 4 kcal mol−1 in the neutral and monocationic species, whereas the stability is reversed in favor of the basal isomer by 3.4 kcal mol−1 in the dicationic species. Considering the energetics of the ligand substitution reaction, ΔGreact decreases from 19 to −2 kcal mol−1 when going from the neutral to the dicationic species, therefore supporting the replacement of a basal CO ligand with CH3CN in the oxidation process. The redox potential calculated in CH3CN for the 2(ap)+ → 2(ba)2+ + e− process (where ap and ba stand for the CH3CN ligand in the apical and basal position, respectively) is equal to 0.56 V, a value about 0.6 V smaller than that calculated for the oxidation of 1+ to 12+ (1.16 V) and that can explain the second oxidation step after the replacement of CO with CH3CN. 2.2. Synthesis, Characterization, and Electrochemistry of [FeMo(CO)4(MeCN)(κ2-dppe)(μ-pdt)] (2). 2.2.1. Synthesis and Characterization of 2. In view to replace carbonyl ligands by acetonitrile in the pentacarbonyl compound [FeMo(CO)5(κ2-dppe)(μ-pdt)] (1), a solution of 1 was heated under reflux, in acetonitrile. Only one product 2 was formed, which could not be gathered as a powder because of its instability in the absence of acetonitrile. The progress of the reaction was followed by IR monitoring, and it was shown that the reaction was conducted to completion after 45 min. The IR spectrum of 2 has four characteristic bands in the ν(CO) region (1942, 1916, 1847, and 1793 cm−1), suggesting the presence of terminal and semi-bridging carbonyl ligands. These ν(CO) bands in 2 are shifted to lower wavenumbers relative to those (2016, 1936, 1882, and 1850 cm−1)36 observed in the spectrum of the parent complex 1, indicating the substitution of carbonyl for a better donor ligand (e.g., CH3CN) (Figure S7, Supporting Information). To characterize compound 2 by NMR spectroscopy, the analyses were carried out in CD3CN solutions of 2, prepared in situ in the NMR tubes (see Experimental Section and Figures S8 and S9, Supporting Information). The characteristic peaks of the coordinated acetonitrile could not be assigned without ambiguity; nevertheless, the 1H NMR spectrum of 2 in CD3CN allowed the assignments of all of the other expected signals associated with the phenyl, dppe, and pdt groups. The 31P{1H} NMR spectrum of 2 in CD3CN shows a singlet at 60.9 ppm assigned to the two phosphorus atoms of the diphosphine, suggesting their equivalence; this can be related to dibasal coordination of the dppe group at the iron atom. The 13C{1H} NMR spectrum in CD3CN shows, in the carbonyl region, the presence of three resonances appearing as a singlet at 225.8, a triplet at 220.6 (2JPC = 17.0 Hz), and a singlet at 219.3 ppm. This pattern is in

a positive shift of the redox potentials (Table 1). The scan rate 1/2 dependence of the current function (ired p /v ) demonstrates that the reduction follows an ECE mechanism at slow scan rates, as in dichloromethane. In contrast, the oxidation of 1 is strongly solvent-dependent. In acetonitrile, the oxidation is irreversible (Figure S5, Supporting Information) and involves more than one electron at moderate and slow scan rates (v ≤ ca. 1 V s−1), while some chemical reversibility is apparent at higher scan rates (v = 10 V s−1, Figure 7). Nevertheless, the

Figure 7. CV of [FeMo(CO)5(κ2-dppe)(μ-pdt)] (1, 1.28 mM) in MeCN−[NBu4][PF6] (vitreous carbon electrode, v = 10 V s−1, potential in V/Fc+/Fc).

value of the cathodic to anodic peak current ratio [(icp/iap)ox = 0.43 for v = 1 V s−1; 0.85 for v = 10 V s−1]69 and the persistence of a reduction peak at −0.12 V on the return scan at fast scan rates (Figure 7) prove that full chemical reversibility is not reached. The substantial increase of the current function of the oxidation at slow scan rates confirms that more than one electron is transferred. From the different nature of the oxidation of 1 in CH2Cl2, a weakly bonding solvent, and in MeCN, a comparatively good ligand, one may conclude that the cation 1+ reacts with MeCN to generate a species easier to oxidize than 1, which results in the transfer of a second electron at moderate and slow scan rates (ECE process). In the case of the diiron analogue, [Fe2(CO)4(κ2-dppe)(μ-pdt)], similar results were obtained and assigned to the binding of MeCN to a vacant site at the rotated cation, followed by further oxidation of the cationic acetonitrile adduct.70 In the present case, such a reactivity is difficult to envisage because the rotation of the {Fe(dppe)CO} entity would cause a steric clash between two semi-bridging carbonyls. The reactivity observed upon oxidation of 1 does not either arise from the simple replacement of the apical CO on the molybdenum center (which the analysis of the X-ray crystal structure suggested to be more labile than the basal ones)36 by MeCN: the CV of the authentic [FeMo(CO)4(MeCN)(κ2-dppe)(μ-pdt)] complex 2 is different from that shown in Figures 7 and S5 (see below). The HOMO of 1 (Figure S6a, Supporting Information), which has a slight Fe−Mo antibonding character, is principally a {Mo-basal CO} bonding orbital. As expected from the nature of the HOMO, the comparison of the DFT-optimized structures of 1 and 1+ shows that oxidation causes a substantial lengthening of the Mo−C(O) bonds of the basal carbonyl ligands (in CH3CN 1: 1.984 Å; 1+: 2.015 Å) and a shortening of the Mo−Fe distance (in CH3CN 1: 2.850 Å; 1+: 2.737 Å). This might indicate that the chemical step of the ECE E


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Inorganic Chemistry

resulting competitive trans influence (e.g., CH3CN) in 2. As expected, the M−C−O (semibridging) angle in 2 [164.8(4)°] is closer to the linearity than that observed in 1 [160.6 (5)°]. The structural parameters of 2, with the exception of the Mo− C(O) semibridging characteristics, are relatively close to those of 1 (Table 4). The DFT-optimized structure of 2 is in very good agreement with the X-ray structure (Figure S10, Supporting Information), with the exception of the Mo1−Fe1 distance (2.839 Å), which is significantly longer than the experimental one. In addition, the isomer featuring dibasal coordination of dppe, as identified by the X-ray analysis, is more stable than the basal−apical isomer by about 8 kcal mol−1. The calculated IR frequencies associated with the CO stretching modes compare fairly well with the experimental ones. In fact, the two frequencies at 1942 and 1916 cm−1 match nearly perfectly the experimental frequencies, whereas the two frequencies at lower wavenumbers (1889 and 1850 cm−1) overestimate the experimental values, as already observed in other dimetallic FeFe clusters.71 The replacement of the apical CO ligand with the CH3CN molecule affects the electronic structure of 2 with respect to 1. The atomic charge of the Mo atom and that of the {Mo(L)4} fragment (−0.63 and −0.37, respectively) are about 0.15 less negative than in 1, indicating a charge transfer from Mo to the {pdt} and the {Fe(dppe)(CO)} fragments of the molecule. This effect can be rationalized by considering an orbital diagram obtained by dividing the cluster into the two neutral {(dppe)(CO)(pdt)Fe} and {Mo(L)4} subunits, which feature the formal FeII and Mo0 redox states, respectively, as done with complex 1 in ref 36 (Figure S6b, Supporting Information). Formation of the metal−metal bond in 2 from the two fragments might be described as a donor−acceptor interaction between the HOMO of {Mo(L)4} and the LUMO of {(dppe)(CO)(pdt)Fe}. Replacement of CO with CH3CN leads to a significant destabilization of the HOMO of the {Mo(L)4} fragment which makes its energy more similar to that of the LUMO of the {(dppe)(CO)(pdt)Fe} fragment. This increases the contribution of the latter orbital in the bonding combination of the HOMO of {Mo(L)4} and the LUMO of {(dppe)(CO)(pdt)Fe}, which corresponds to the HOMO − 3 orbital in 2. This orbital, which is characterized by a shape suitable to bind H+ to give the bridging hydride, is also shifted to an energy higher than in 1 by about 0.4 eV suggesting that 2 is more basic than 1. 2.2.2. Electrochemistry of [FeMo(CO)4(MeCN)(κ2-dppe)(μpdt)] (2) in MeCN−[NBu4][PF6]. The electrochemical behavior of 2 has been investigated only in MeCN because this complex is not stable in other solvents (see Experimental Section). The scan rate dependence of the oxidation peak current function 1/2 (iox p /v ) and the CV of the complex (Figures S11 and S12, Supporting Information) demonstrate that 2 undergoes a reversible one-electron oxidation at E1/2 = −0.13 V. While the oxidation of 1 in MeCN showed some chemical reversibility only at fast scan rates, that of 2 appears to be reversible at the slowest scan rate used in our study (v = 0.05 V s−1). The replacement of the apical CO on the Mo center of 1 by an

agreement with a {Fe(CO)P2}(μ-CO){Mo(CO)2L} core, by comparing the spectrum of 2 with that of the parent complex 1.36 Therefore, these spectroscopic results suggest the substitution of the apical CO coordinated to molybdenum for an acetonitrile, as depicted in Scheme 2. Indeed, the Mo− Scheme 2. Synthesis of 2

C(O) apical distance (2.085 (7) Å) in 1 is significantly longer than those of the equatorial CO ligands (average = 1.991 Å), so this accords with a higher reactivity based on the lability of this carbonyl group.36 The formulation of complex 2 was confirmed by X-ray analysis of a single crystal of 2·2CH3CN (Figure 8), obtained

Figure 8. ORTEP of [FeMo(MeCN)(CO)4(κ2-dppe)(μ-pdt)] (2) in crystals of 2·2CH3CN. Here, non-hydrogen atoms are shown with ellipsoids at the 30% probability level. H atoms bonded to C atoms are omitted for clarity. Selected distances (Å) and angles (deg) are as follows: Mo1−Fe1 = 2.7645 (7), Mo1−C2 = 1.946 (5), Mo1−C4 = 1.969 (5), Mo1−N40 = 2.206 (4), Fe1−C1 = 1.735 (5), Fe1−C2 = 2.581, N40−C40 = 1.149 (7), C1−O1 = 1.152 (6), C2−O2 = 1.172 (5), C3−O3 = 1.162 (6), C4−O4 = 1.156 (6), Mo1−C2−O2 = 174.2, Mo1−C3−O3 = 174.2 (4), Mo1−C4−O4 = 178.1 (4), Fe1− C1−O1 = 176.3 (4), Mo1−N40−C40 = 173.2.

from a mixture of acetonitrile/diethyl ether at low temperature. In 2, the molybdenum atom lies in a pseudo-octahedral geometry, while the iron atom is in a typical square-pyramidal geometry supplemented by an interaction with a semibridging carbonyl [Mo1−C2, 1.946(5) Å; Fe1−C2, 2.581 Å; Mo1− C2−O2, 164.8(4)°]. The Mo1−C2(O) distance is significantly shorter than the Mo−C(O) terminal one (∼1.967 Å), in accordance with the substitution of the apical CO in 1 and the

Table 4. Comparison between 1 and 2 Selected Structural Parameters (Å, deg) compound

Mo1−Fe1

Mo1−C2

Mo1−COt(basal)

Fe1−C2

C2−O2

average C−Oterminal

Mo1−C2−O2

Mo1−C4−O4

136 2

2.7795(8) 2.7645(7)

1.987(6) 1.946(5)

1.984 1.967

2.514 2.581

1.173(7) 1.172(5)

1.13 1.16

160.6(5) 164.8(4)

175.6(7) 178.1(4)

F


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Figure 9. Multiple scan CV of [FeMo(CO)4(MeCN)(κ2-dppe)(μ-pdt)] (2; ca. 1 mM) in MeCN−[NBu4][PF6]; the marked peak is probably due to the formation of 1 as a minor byproduct; see text (vitreous carbon electrode, potential in V/Fc+/Fc).

arrow in Figure 9b might be due to the intermediate P2, it might also arise from the formation of complex 1 as a minor side-product because it occurs at the same potential as the reduction of [FeMo(CO)5(κ2-dppe)(μ-pdt)] (1). In this case, the oxidation of [P2]− would produce directly the neutral complex 2, and the coupling of the chemical reaction (MeCN binding) to the electron transfer may explain the highly irreversible nature of the latter. In contrast, the oxidation of [P2]2− is a reversible one-electron process (Scheme 3) as clearly shown by the comparison of the peak current of the corresponding system under steady-state conditions (Figure S13, red trace) to that of the reversible one-electron oxidation of complex 2 (Figure S13). The nature of the chemical step following the reduction of 2, that is, the loss of the acetonitrile ligand, is inferred from the comparison of the cyclic voltammograms of 2 under N2 and under CO (Figure S14, Supporting Information). When CO is present, the oxidation peak of [P2]2− (Eox1 1/2 = −1.86 V) is not observed and that of [FeMo(CO)5(κ2-dppe)(μ-pdt)]− (Eox 1/2 = −1.70 V, Table 1)75 is detected instead, showing that the reaction of [P2]− with CO generates the all-carbonylated anion 1−. The lack of stability of 2 in the absence of an excess of acetonitrile suggested that this ligand is labile in the neutral molecule (see above). The electrochemical results show that the lability of MeCN is strongly enhanced by the reduction of the complex. An approximate value of the rate constant of the loss of MeCN was obtained from the current functions in Figure S12, using the method of Nicholson and Shain76 adapted by Geiger77 (see Supporting Information for details). Digital simulations using the estimated value of the rate constant of the MeCN loss, kc ≈ 1280 s−1 (Figure S15, Supporting Information), produce cyclic voltammograms in coherence with the experimental ones (Figure S16, Supporting Information). The calculated oxidation potential for the 2 → 2+ process in acetonitrile is equal to −0.21 V (exp. = −0.13 V), a value about 0.35 V lower than that calculated for the 1 → 1+ process, and in good agreement with the experimental value. The reduction of 2 to give 2− occurs at −2.06 V, a value slightly more negative than that calculated for 1, and also in this case, in good agreement with the experiment (−2.03 V). To investigate the reactivity of the reduced species 2−, we have considered dissociation energies of CH3CN to the Mo atom as a function of the charge of the complex. As shown in Table 4, ligand dissociation energies increase by augmenting the negative charge of the complex, indicating that CH3CN is more labile after the reduction of the neutral complex.

acetonitrile molecule thus shifts the oxidation of the complex toward more negative potentials (by ca. 0.3 V, Table 1), as expected, and improves noticeably the chemical stability of the cation. The comparison of the current functions (Figure S12, Supporting Information) demonstrates that the reduction of ox complex 2 involves two electrons (ired p /ip = 2.08 for v = 0.05 V s−1; 2.23 for v = 1 V s−1; 1.71 for v = 50 V s−1) and leads to a product detected on the reverse scan by two oxidation processes. The CVs in Figure 9 show two important pieces of information. First, despite the apparent partial reversibility of the reduction, the oxidation peak observed on the return scan around −1.8 V belongs to another system. Indeed, the curve crossings in Figure 9a show that the reduction produces a species whose oxidation occurs at a potential (E1/2 = −1.86 V; also see red trace in Figure S13) very close to that of the reduction of 2 (Ered p = −2.03 V at the scan rate used for Figure 9). Second, the second oxidation of the reduction product, characterized by a very broad peak (Eox p ≈ −1.2 V), is a highly irreversible process which partially regenerates 2 (Figure 9b). This is evidenced by the comparison of the reduction currents of the successive scans in Figure 9a to those in Figure 9b: the reduction current at ca. −2 V is slightly larger when the oxidation around −1.2 V is included in the potential range than when it is not. Therefore, 2 is reduced following an ECE mechanism to a product ([P2]2−, Scheme 3) whose oxidation regenerates 2 according to an EEC process. The association of a forward ECE to a backward EEC mechanism has already been observed.72−74 The neutral P2 is shown between braces in Scheme 3 because it might be a short-lived, undetected intermediate. Indeed, although the small reduction peak marked with an Scheme 3. Proposed Mechanisms for the Reduction of [FeMo(CO)4(MeCN)(κ2-dppe)(μ-pdt)] (2) and for the Oxidation of the Product [P2]2− in MeCN−[NBu4][PF6]

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exergonic by about 2 kcal mol−1 (Table S4, Supporting Information). However, dissociation of the phosphine ligand should be prevented by dissociation of CH3CN with the subsequent further reduction of the [P2]− monoanion. 2.3. Electrochemistry of [FeMo(CO)4(L)(κ2-dppe)(μpdt)] in the Presence of Acid (L = CO (1), MeCN (2)). 2.3.1. [FeMo(CO)5(κ2-dppe)(μ-pdt)]. The reactivity of complex 1 toward acids of moderate strength (CH3SO3H, pKa = 10,78 8.4 in MeCN;79 HOTs, pKa = 8.6 in MeCN)80 was investigated in MeCN (Figure 10) and in CH2Cl2 (Figure S18, Supporting Information). The cyclic voltammograms recorded under such mildly acidic conditions show that the oxidation of the complex is almost unaffected (Figures 10a and S18a), which proves that 1 is not protonated by these acids. The new reduction peak observed at −1.6 V thus arises from the reduction of 1 that is shifted positively because of the protonation of 1− to 1H. The increase of the reduction peak current at the lowest acid concentrations indicates that 1H is in turn reduced at a potential equal to or less negative than −1.6 V (ECE mechanism, see Section 2.4 below). The reduction peak current continues to increase upon further additions of acid (Figures 10b and S18b), indicating that a catalytic proton reduction process takes place. However, the current around −1.6 V culminates, after addition of ca. 30−35 equiv of acid, to a value less than 10 times larger than the reduction current measured in the absence of acid (Figure 10b), which implies that the catalysis is probably limited by the slow release of dihydrogen (see Section 2.4 below). When an excess of the stronger acid CF3SO3H (pKa = 0.7 in MeCN, pKa = −11.4 in 1,2-dichloroethane)81,82 is added to a dichloromethane solution of complex 1,83 the acid−base equilibrium in Scheme 4 is shifted to the right.

Dissociation of the CH3CN ligand in the monoanion might also be assisted by an isomerization process in which the terminal CO on Fe moves to a bridging position between the two metal atoms, creating a vacant coordination site on the Fe atom (Figure S17, Supporting Information). This isomer ([P2]iso2− in Table 4 and Figure S17) is more stable than the one in which the free coordination site is on Mo ([P2]iso1− in Table 5 and Figure S17) by about 9 kcal mol−1 and we can Table 5. Dissociation Gibbs Free Energies (ΔGdiss, kcal mol−1) of the Apical (Mo···CH3CN) Ligand in 2, 2−, and 22−, and Relative Stabilities (ΔGSO, kcal mol−1) of the iso1 and iso2 Isomers Calculated in Vacuum and CH3CN Solvent ΔGdiss 2 → {P2}iso1 + CH3CN 2 → {P2}iso2 + CH3CN 2− → [P2]iso1− + CH3CN 2− → [P2]iso2− + CH3CN 22− → [P2]iso12− + CH3CN 22− → [P2]iso22− + CH3CN

ΔGSO

vacuum

CH3CN

vacuum

CH3CN

3.5 12.9 −9.8 −13.5 −27.5 −28.8

0.3 8.4 −11.9 −21.0 −37.4 −37.0

0.0 9.3 3.8 0.0 1.3 0.0

0.0 8.1 9.1 0.0 0.0 0.4

assign the [P2]− species detected by the experiments to this form. Considering this rearrangement, the dissociation of the CH3CN ligand in the monoanion is a spontaneous process by about 21 kcal mol−1. In the case of the dianion dissociation of CH3CN is still more favorable, with dissociation Gibbs free energy equal to about −37 kcal mol−1. A further structural rearrangement occurs in [P2]iso22− because in the most stable isomer ([P2]iso12− see Figure S17), the dppe prefers the basal−apical arrangement, and all CO groups move to a terminal position. It is worth noting that the second reduction from [P2]iso2− to [P2]iso12− occurs at a calculated potential of −2.03 V, a value very similar to the one measured from the experiment, and that can explain the dielectronic reduction observed in the cyclic voltammogram. On the other hand, reduction of 2− to give 22− should occur at a much more negative potential calculated at −2.74 V. For the reduction of 2, we have also investigated a mechanism similar to that discussed for 1, in order to disclose the factors discriminating between the two different processes. It is worth noting that dissociation of the phosphine ligand can still occur in the monoanion, as the isomerization process is

Scheme 4. Oxidation of [FeMo(CO)5(κ2-dppe)(μ-pdt)] (1) in Acidic Medium (HA = CF3SO3H)

The occurrence of the protonation of 1 by a strong acid into the bridging hydride (see Section 2.4 below) is evidenced by the color change of the solution from dark green to orange, and by the decay of the complex oxidation peak current in CV.

Figure 10. CV of [FeMo(CO)5(κ2-dppe)(μ-pdt)] (1, 0.82 mM) in MeCN−[NBu4][PF6] in the presence of increasing amounts of HOTs (vitreous carbon electrode, v = 0.2 V s−1, potential in V/Fc+/Fc). H


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sharp contrast with the reduction of the diiron analogue, [Fe2(CO)4(κ2-dppe)(μ-pdt)(μ-H)]+, which was shown to catalyze the reduction of strong acids at ca. −1.3 V.84,87 2.3.2. [FeMo(CO)4(MeCN)(κ2-dppe)(μ-pdt)]. The replacement of an electron-withdrawing CO ligand (ligand parameters PL = 0 V;88 EL = 0.99 V89) by the comparatively good donor ligand MeCN (ligand parameters PL = −0.58 V;88 EL = 0.34 V89) is expected to increase the basicity of the complex and thus to facilitate its protonation by a given acid (Scheme 4). That it is the case is shown by the fact that the decrease of the oxidation peak of the complex [FeMo(CO)4(L)(κ2-dppe)(μpdt)] upon additions of CF3SO3H is faster when L = MeCN than when L = CO (Figures 11a and S19a, Supporting Information). The addition of CF3SO3H to acetonitrile solutions of 2 thus leads to the formation of a hydride complex. In addition, a new partially reversible system is detected at E1/2 = −0.41 V (Figure S19a) and shown to correspond to a reduction process (Figure S19b). This suggests that two different sites might be accessible to protons in complex 2: reaction at the first one generates the protonated complex 2H+, whereas protonation at the other one leads to an unknown oxidized species (and presumably 1/2H2). As for the precursor 1, the reduction of the hydride of 2 at −1.14 V in the presence of acid does not result in any catalytic process. Catalysis of proton reduction is observed at a more negative potential (E pred ≈ −1.68 V, Figure S20, Supporting Information), but this poorly efficient process was not investigated further. 2.4. Spectroscopic and Theoretical Analyses of the protonation of [FeMo(CO)4(L)(κ2-dppe)(μ-pdt)] [L = CO (1), MeCN (2)]. The protonated species [FeMo(CO)5(κ2dppe)(μ-pdt)(μ-H)]+ (1H+) and [FeMo(MeCN)(CO)4(κ2dppe)(μ-pdt)(μ-H)]+ (2H+) were prepared in situ in a NMR tube by addition of 3 or 2.5 equiv of HBF4·Et2O to a CD2Cl2 or CD3CN solution of 1 or 2, respectively (see Experimental Section). In both cases, the solution rapidly turned from darkgreen to orange (1) or red-orange (2), indicating the formation of new complexes. Then, the solutions were analyzed by IR and NMR techniques. The two compounds have not been isolated in the solid state because of their low stability. Several attempts to crystallize 1H+ and 2H+ in a mixture of CH2Cl2/diethyl ether, at −80 °C, failed. Therefore, these compounds have been only characterized by their spectroscopic data. The IR spectra of 1H+ and 2H+ show four or three characteristic bands in the ν(CO) region, respectively (see the Experimental Section). These ν (CO) bands are significantly shifted to higher wavelengths relative to those of the parent complexes, which supports the formation of cationic species. NMR recording reveals a complete and clean transformation, at low temperature, of 1 into a new species 1H+. The 31P{1H} pattern of 1H+, at −40 °C, is characterized by a signal at 79.05 ppm. The 1H NMR spectrum (−40 °C) reveals the presence of a triplet at −11.11 ppm with a coupling constant 2JPH of 22 Hz that is fully consistent with the formation a bridging hydride species with a diphosphine in a dibasal conformation. 13 C NMR spectra recorded at this temperature reveal, in the carbonyl region, the presence of a singlet at 216.7 ppm that could be assigned to the carbonyl bound on the molybdenum atom, and a triplet at 213.8 ppm with a coupling constant 2JPC of 20 Hz that can be unambiguously attributed to the carbonyl linked to the {FeP2} moiety. This pattern strongly suggests that the protonated species 1H+ does not feature a semi-

From Figure 11a, it can be seen that the oxidation of 1 in the presence of an excess of CF3SO3H takes place according to a

Figure 11. CV of [FeMo(CO)5(κ2-dppe)(μ-pdt)] (1, 1.03 mM) in CH2Cl2−[NBu4][PF6] in the presence of increasing amounts of CF3SO3H (vitreous carbon electrode, v = 0.05 V s−1, potential in V/ Fc+/Fc).

CE mechanism.38−40 The S-shape of the positive scan (whereas the reverse scan is still peak-shaped) is characteristic of a process in which the (oxidation) current is controlled by the kinetics of the preceding equilibrium, here the protonation of complex 1 (Scheme 4), rather than by the diffusion of 1 toward the electrode. The capability to shift the protonation equilibrium to the left at the electrode surface suggests that this equilibrium is mobile and that the pKa’s of the protonated complex 1μH+ and of CF3SO3H are not very different. The protonation of 1 is signaled by a new peak at −1.04 V in the cyclic voltammogram (Figure 11b). The 0.8 V difference between this reduction peak and that of the neutral complex 1 is consistent with the formation of a bridging hydride (Scheme 4),84−86 which has been established by the spectroscopic analyses of the protonated product, as well as by theoretical calculations (see Section 2.4 below). The ratio of the reduction current measured at −1.04 V in the presence of a large excess of acid (40 equiv, Figure 11b) to the oxidation current measured in the absence of acid shows that no more than two electrons are involved in the reduction. In the hypothesis of an ECE reduction mechanism at −1.04 V, the fact that no catalysis takes place at this potential might indicate that the chemical step of the ECE process is not the protonation of 1μH. Alternatively, if one assumes that 1μH is protonated and the product further reduced at a potential equal to (or less negative than) −1.04 V, the absence of catalytic current would demonstrate that no dihydrogen is released at all. This is in I


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Inorganic Chemistry bridging carbonyl ligand (Figure 12a). Similar IR, 1H, and 31P NMR data were obtained for 2H+, which are also consistent

the acidity of the two complexes can be explained by considering the asymmetric charge distribution in 1 which makes the Fe atom “partially oxidized” to the Fe(II) redox state as already discussed.36 Indeed, the atomic charge of Fe in 1 is less negative than in 1Fe. In addition, the three highest occupied orbitals of 1 are mainly localized on Mo, and only the HOMO − 3 orbital is suitable to interact with H+ to give the bridging hydride, whereas in 1Fe, the HOMO can be used for the interaction with the H+ ion (Figure S5a, Supporting Information). It is interesting to note that protonation of 1− leading to 1H is complete also with the weak CH3SO3H acid (Kc ≈ 103), supporting the experimental observation that protonation occurs after the one-electron reduction of 1. However, as stated in the previous section, 1− can rearrange by dissociating a PPh2 arm of the diphosphine ligand with the concomitant migration of the semibridging CO from Mo to the Fe atom, yielding the P− species. Protonation of either 1− or P− can then proceed by two distinct mechanisms (Scheme 5).

Figure 12. (a) Proposed structure for 1H+ on the basis of spectroscopic analysis and (b) DFT-calculated geometry of the most stable 1H+ isomer in CH3CN.

with the formation of a μ-hydride species, and will not be more commented on (see the Experimental Section). An extensive DFT search on the PES of 1H+ identified the μ-hydride species as the most stable isomer. The geometry of such species, calculated in the CH3CN implicit solvent, is shown in Figure 12b. The hydrogen atom asymmetrically bridges the two metals, being slightly closer to Fe (1.682 Å) than Mo (1.937 Å). The semibridged CO moves to a terminal position on Mo, which therefore is heptacoordinated with a capped trigonal prismatic arrangement of ligands. The Fe−Mo distance (2.849 Å) is very similar to that of the parent compound 1. The calculated IR frequencies at 1969, 1999, 2004, 2009, and 2075 cm−1 fairly agree with the experimental spectrum by assigning the very close calculated frequencies at 1999, 2004, and 2009 cm−1 to the peak centered at 2025 cm−1. DFT calculations confirm that 1 is much less basic than the corresponding homonuclear [Fe2(CO)4(κ2-dppe)(μ-pdt)] cluster (1Fe). Indeed, calculation of the equilibrium constants reported in Table 6 shows that protonation of 1 does not occur with the weak acid CH3SO3H, (Kc ≈ 10−13) while it can partially occur with the strong CF3SO3H acid (Kc ≈ 10−4). On the contrary, the protonation of the diiron analogue [Fe2(CO)4(κ2-dppe)(μ-pdt)] is already complete in the presence of the weak CH3SO3H acid. The large difference in

Scheme 5. Proposed Concurrent Mechanisms for the Formation of H2 from 1 with the Weak CH3SO3H Acid

Protonation of 1− yields 1H (Scheme 5, Path Ablack), in which the hydride ion is terminally coordinated to the Fe atom (Fe−H = 1.576 Å; Mo−H = 2.387 Å), and the Fe−Mo distance is equal to 3.112 Å leaving the Mo atom with an octahedral coordination (Figure S21, Supporting Information). The dramatic increase in the basicity of 1− with respect to 1 can be explained by considering the electronic structure of the anion. Indeed, as discussed in the previous section, reduction of 1 occurs at the Fe atom, which, therefore, changes from the Fe(II) to the Fe(I) redox state. In this respect, it is worth noting that the basicity of 1− (evaluated by considering the equilibrium constants reported in Table 6) is similar to that of the 1Fe complex featuring a “genuine” Fe(I) redox state. The redox potential of the 1H + 1e− → 1H− process calculated in CH3CN is equal to −0.89 V, a value more positive than that calculated for the 1 + 1e− → 1− process, supporting the second reduction after protonation of 1−. In 1H−, the Fe−Mo distance further increases up to 3.688 Å, and the hydride ion moves to a full terminal position on the Fe atom (Figure S21). Protonation of 1H− at the hydride ion to give the H2 adduct (1H2) is complete with both CH3SO3H, (Kc ≈ 104) and CF3SO3H (Kc ≈ 1014) acids. The release of H2 from 1H2 is a very exergonic step as the dissociation energy of H2 to give the parent complex 1 is equal to about −25 kcal mol−1. Considering the mechanism described above 1 should efficiently catalyze the reduction of protons at the potential of

Table 6. Equilibrium Constant Kc Calculated for the Reaction Y + AH → YH+ + A− (AH = CH3SO3H, CF3SO3H) in CH3CNa AH = CH3SO3H 1Fe + AH → 1FeH+ + A− 1FeH + AH → 1FeH2+ + A− 1 + AH → 1H+ + A− 1− + AH → 1H + A− 1H + AH → 1H2+ + A− 1H− + AH → 1H2 + A− P− + AH → PH + A− PH + AH → P(H)(H)+ + A− PH− + AH → P(H)(H) + A− 1H + AH → P(HS)(H)+ + A− 1H− + AH → P(HS)(H) + A− 2 + AH → 2H+ + A− 2− + AH → 2H + A− 2H− + AH → 2H2 + A− 2PH− + AH → 2PH2 + A−

1.0 8.4 8.0 2.5 1.1 2.2 1.4 2.4 1.4 6.4 3.6 3.4 1.3 6.9 1.3

× × × × × × × × × × × × × × ×

102 10−8 10−14 103 10−11 104 104 10−18 10−1 10−8 109 10−6 109 106 105

AH = CF3SO3H 3.2 2.7 2.5 6.4 3.4 7.0 4.3 7.7 1.5 2.0 1.1 1.1 4.1 2.2 4.1

× × × × × × × × × × × × × × ×

1011 102 10−4 1012 10−2 1013 1013 10−9 108 102 1019 104 1018 1016 1014

a

Details of calculations are reported in the Experimental Section. J


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Inorganic Chemistry −1.6 V. However, the isomerization of 1− to P−, as discussed in the previous section, can change the mechanism (Scheme 5, Path Bred). Indeed, P− can be protonated at the bridging position yielding PH (Figure S21), which is about 5 kcal mol−1 more stable than 1H. The redox potential of the PH + 1e− → PH− process is equal to −0.48 V, and the reduced species is as much as 14.7 kcal mol−1 lower in energy than 1H−. Protonation of PH− occurs at the Mo atom yielding the dihydride P(H)(H) species (Figure S21) which is about 8 kcal mol−1 more stable than 1H2. The H2-adduct (PH2) is not a stable species as it spontaneously evolves during geometry optimization to P(H)(H). Notably, dissociation of H2 from P(H)(H) leading to P, is an endergonic process by 5.4 kcal mol−1. Therefore, the release of H2 through this pathway should be prevented by the formation of the stable P(H)(H) intermediate from which the dissociation of H2 is an unfavorable process. The isomerization of P to 1, which is exergonic by 23 kcal mol−1, restores the starting complex. The two concurrent mechanisms for the reduction and protonation of 1 by weak acids are summarized in Scheme 5. As stated above, protonation with the strong acid can also occur at the neutral complex to give the 1H+ species, which is then reduced to 1H. The calculated redox potential of the 1H+ + 1e− → 1H process in CH3CN is equal to −0.82 V, in fairly good agreement with the experimental value of −1.04 V. The protonation of 1H, also considering the strong CF3SO3H acid, is not completely shifted toward the products, as the corresponding Kc is ∼10−2. Again, the corresponding diiron complex [Fe2(CO)4(H)(κ2-dppe)(μ-pdt)] is significantly more basic because the protonation reaction is complete also with the weak CH3SO3H acid. This difference in basicity may explain the lack of catalytic behavior of the FeMo complex in these conditions; the small value of the equilibrium constant should favor the isomerization reaction of 1H to the more stable species PH, that is then reduced at the potential of −0.48 V and protonated to give the P(H)(H) product. It is worth noting that also protonation of 1H does not lead to the H2 adduct 1H2+, but it occurs at the sulphur atom of the dithiolate ligand, with the concomitant cleavage of the S−Mo bond, yielding the P(HS)(H)+ species (Figure S21). Indeed, this latter species is more stable than the H2-adduct (1H2+) by about 7 kcal mol−1, and the corresponding Kc is larger by about four orders of magnitude (Table 6). Reduction of such species leads to P(HS)(H) (Figure S21) which is still more stable than 1H2 by about 5 kcal mol−1. Therefore, both pathways lead to dead-end products from which H2 is not released. 2 is significantly more basic than 1, as confirmed by the Kc associated to the proton transfer reaction 2 + CX3SO3H (X = H, F) → 2H+ + CX3SO3− (X = H, F) which is about 6 orders of magnitude larger than that calculated for 1 (Table 6). However, the value calculated for the reaction with the weak CH3SO3H acid is still very small (Kc ≈ 10−6), indicating that the equilibrium is almost completely shifted toward the reactants, whereas in the case of the CF3SO3H acid 2 is fully protonated (Kc ≈ 104) to give the hydride 2H+ as observed experimentally. The geometry of 2H+, calculated in CH3CN, is similar to that of 1H+; the bridging hydrogen atom is slightly closer to Fe (1.704 Å) than Mo (1.888 Å), and the Fe−Mo distance is 2.867 Å (Figure S22, Supporting Information). The catalytic production of H2 from 2H+ should occur according to an ECE mechanism initiated by its reduction to 2H, followed by protonation of the hydride ion to give the H2-adduct and the second reduction for the release of H2. In fact, the

calculated redox potential of the 2H+ + 1e− → 2H process (−1.20 V) is in perfect agreement with the experiment and slightly lower than that calculated for 1H+. In 2H, the hydride ion moves closer to the Fe atom (Fe−H = 1.625 Å; Mo−H = 2.093 Å), although to a lesser extent than in 1H, whereas the Fe−Mo distance increases up to 3.073 Å (Figure S22). Protonation of 2H yields the 2H2+ species, for which the H2adduct, with H2 coordinated to the Fe atom, has been identified as the most stable isomer (Figure S22). The Kc associated with protonation of 2H by CH3SO3H and CF3SO3H is equal to about 10−9 and 1, respectively, indicating that such a process is favored in the presence of the strong acid. Reduction of 2H2+ at the calculated redox potential of −0.53 V leads to 2H2 (Figure S17), in which dissociation of H2 is an exergonic process by as much as 27.2 kcal mol−1. According to the mechanism discussed above, 2 should efficiently catalyze the production of H2 at the potential of −1.20 V, in contrast with the experimental evidence. However, as observed for 1, the mechanism may be modified by the formation of more stable intermediates along the reaction pathway. Indeed, we have found that release of the CH3CN ligand can easily occur after the formation of 2H. Dissociation of CH3CN from this species is an exergonic process by as much as −33 kcal mol−1 and leads to the P2H intermediate. Interestingly, P2H can be reduced at a calculated redox potential of −0.78 V, a value more positive than that calculated for the reduction of 2H+, indicating that the second reduction can occur after the formation of P2H to give the P2H− intermediate. Notably, protonation of such species does not occur at the hydride ion to give the H2-adduct, but at the Mo atom forming the dihydride species P2(H)(H) (Figure S17). Indeed, the H2-adduct was not isolated as a minimum on the PES as it spontaneously evolved during geometry optimization to P2(H)(H). Such protonation is favored when P2H− is treated with both CH3SO3H and CF3SO3H; the corresponding Kc values are equal to about 105 and 1014, respectively. Release of H2 from P2(H)(H) is an endergonic process by only 1 kcal mol−1, but this energy term does not consider the formation of H2 on the complex, which is not an energy minimum on the PES. In addition, such ΔG includes a significant positive entropic contribution due to the dissociative nature of the reaction as, for example, the ΔH of the same reaction is more than 12 kcal mol−1. In summary, our results suggest that dissociation of CH3CN from 2H leads to a species (P2H) that further evolves to a dihydride dead-end product from which H2 is not formed, explaining the lack of catalytic activity of this MoFe cluster.

3. CONCLUSIONS The cooperativity of metal centers in polymetallic systems is key to the efficiency of the active sites of several metalloenzymes. In this work, we have explored by CV and DFT studies the behaviors of heterobimetallic species featuring a FeMoS2 framework with connected iron and molybdenum centers, as truncated models of the FeMo-co. Unfortunately, dinitrogen does not coordinate at this FeMoS2 core despite the fact that a coordination site is free at the molybdenum atom, as shown by the formation of the MeCN-adduct 2. It is worth noting that in this species, an isomerization process through a CO-migration is proposed to favor the dissociation of the acetonitrile upon reduction (Scheme 6). Calculations suggest the possibility to switch the coordination site from one metal center to the other one. The K


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Inorganic Chemistry

Data for 2 are as follows. IR (CH3CN, cm−1): ν(CO) 1942 (m), 1916 (s), 1847 (s), 1793 (m). 1H NMR (CD3CN): δ 7.88−7.19 (m, 20H, Ph), 3.08 (m, 2H, P(CH2)2P), 2.72 (m, 2H, S(CH2)3S), 2.68 (m, 2H, P(CH2)2P), 2.53 (m, 2H, S(CH2)3S), 2.26 (m, 2H, S(CH2)3S). 31P{1H} NMR (CD3CN): δ 60.9(s). 13C{1H} NMR (CD3CN): δ 225.8 (s, μ-CO), 220.6 (t, 2JPC = 17.0 Hz, Fe(CO)), 219.3 (s, Mo(CO)basal), 138.4 (m, 1JPC = 44.0 Hz, CipsoPh), 136.0 (m, 1 JPC = 27.7 Hz, CipsoPh), 132.8 (pt, CmetaPh), 132.4 (pt, CmetaPh), 130.5 (s, CparaPh), 130.3 (s, CparaPh), 129.2 (pt, CorthoPh), 129.1 (pt, CorthoPh), 31.5 (s, SCH2CH2CH2S), 25.7 (dd, 1JPC = 22.6 Hz, 2JPC = 18.9 Hz, P(CH2)2P), 24.5 (s, SCH2CH2CH2S) (coordinated acetonitrile resonances are not assigned). 4.2.2. Protonation of [FeMo(CO)5(κ2-dppe)(μ-pdt)] (1). 1 (8.7 mg, 0.11 mmol) was purified by filtration under celite, which was solubilized in CD2Cl2 in a NMR tube. HBF4·Et2O (4 μL, 0.033 mmol, 3 equiv) was added under nitrogen, at low temperature. Rapidly, the solution turned from dark-green to orange and then was analyzed by IR and NMR spectroscopy. Data for 1H+ are as follows. IR (−30 °C, cm−1): ν(CO) 2097 (m), 2050 (sh), 2025 (s), 1972 (s). 1H NMR (CD2Cl2, −40 °C): δ 7.72− 7.06 (m, 20H, Ph), 3.26 (m, 2H, P(CH2)2P), 2.95 (2× m, 2H + 4H, S(CH2)3S + P(CH2)2P), 2.80 (m, 2H, S(CH2)3S), −11.11 (t, 2JPH = 22.5 Hz, μ-H). Two protons of the dithiolate bridge (S(CH2)3S) could not be assigned unambiguously, and we assumed that they are hidden by the signals of Et2O (HBF4·Et2O). 31P{1H} NMR (CD2Cl2, −40 °C): δ 79.0 (s). 13C NMR (J-MOD) (CD2Cl2, −40 °C): δ 216.7 (s, Mo(CO)4), 213.8 (t, 2JPC = 20 Hz, Fe(CO)), CipsoPh were not detected; {132.1, 131.7,131.3, 129.7, 129.1, Ph}, 36.7 (s, SCH2CH2CH2S), 26.9 (P(CH2)2P), 23.5 (s, SCH2CH2CH2S). 4.2.3. Protonation of [FeMo(CO)4(MeCN)(κ2-dppe)(μ-pdt)] (2). To a solution of 2 (0.045 mmol) in acetonitrile (10 mL) was added 2.5 equiv of HBF4·Et2O (0.11 mmol, 15.4 μL) at low temperature. Rapidly, the solution turned from dark-green to red-orange, indicating the formation of a new product 2H+, which was analyzed by IR. In order to characterize the unstable protonated adduct 2H+ by NMR spectroscopy, this complex was prepared in situ in an NMR tube by the addition of 4 equiv of CF3SO3H (0.062 mmol, 5.5 μL) to a CD3CN solution of 2 (ca. 0.015 mmol), prepared also in situ; immediately, the solution turned from green to red and then was analyzed. Data for 2H+ are as follows. IR (CH3CN, cm−1): ν(CO) 2049 (s), 1975 (s), 1954 (s). 1H NMR (CD3CN): δ 7.72−7.06 (m, 20H, Ph), 3.2 (m, 2H, P(CH2)2P), 2.9 (m, 2H, P(CH2)2P), 2.3 (m, 1H, S(CH2)3S), 2.6 (m, 2H, S(CH2)3S), −12.81 (t, 2JPH = 18.5 Hz, μ-H). 31 1 P{ H} NMR (CD3CN): δ 76.7 (s). 4.3. X-ray Structural Determination. Measurements for compound 2·2CH3CN were carried out on a Oxford Diffraction XCalibur-2 CDD diffractometer equipped with a jet cooler device. Graphite-monochromated Mo Kα (λ = 0.71073 Å) was used in this experiment. The structure was solved and refined by standard procedures.89,90 Selected bond lengths and angles are mentioned in the caption of Figure 8. Crystal data collection and processing are given in Table 7. 4.4. DFT Calculations. Geometry optimizations and energy calculations have been carried out in the DFT framework with the TURBOMOLE 7.01 suite of programs91 by using the BP86 functional91,92 in conjunction with the resolution-of-the-identity technique93 and a valence triple-ζ basis set with polarization functions on all atoms (TZVP).94 For the Mo atom, a pseudopotential for the 28 inner electrons has been used.95 Stationary points of the energy hypersurface have been located by means of energy gradient techniques, and full vibrational analysis has been carried out on the structures optimized in vacuum to further characterize each stationary point. Free energy (G) values have been obtained from the electronic SCF energy considering three contributions to the total partition function (Q), namely, qtranslational, qrotational, and qvibrational, under the assumption that Q may be written as the product of such terms.96 Enthalpy and entropy contributions have been evaluated by considering the values of temperature, pressure, and scaling factor

Scheme 6. Proposed Isomerisation Process in {FeMo} Species after Reduction and Release of a Substrate (Herein, MeCN)

fact that these FeMo complexes are poorer electrocatalysts for the reduction of proton into dihydrogen than their counterparts may remind that the FeMo nitrogenase act as a better nitrogenase and a poorer hydrogenase that the Fe−Fe analogue. Finally, these results raise also the question of the coordination of N2 at a metal−sulfur site. Studies are now under scope for building more efficient FeMo platforms for N2 activation, using amine-functionalized-diphosphine and limiting the number of carbonyl groups at the heterobimetallic core also in order to increase its basicity.

4. EXPERIMENTAL SECTION 4.1. General. All the experiments were carried out under dinitrogen or argon, using standard Schlenk techniques for the syntheses. All reagents were used as purchased (Sigma-Aldrich). The solvents were predried using conventional methods and were distilled prior to use. The complex [FeMo(CO)5(κ2-dppe)(μ-pdt)] (1) was obtained following the reported procedure.36 The NMR spectra (1H, 31 P, and 13C) were recorded with Bruker AMX 400, AC300, and DRX 500 spectrometers, respectively, of the “Service de RMN de l’Université de Bretagne Occidentale, Brest”. 2D NMR and variable temperature experiments were carried out on a Bruker DRX 500 spectrometer. The infrared spectra were recorded on a FT IR VERTEX 70 Bruker spectrometer. Chemical analyses were made by the “Service de Microanalyse I.C.S.N.”, Gif sur Yvette (France), or the “Centre de Microanalyse du CNRS”, Vernaison (France). 4.2. Electrochemistry. The preparation and the purification of the supporting electrolyte [NBu4][PF6] were as described previously.41 The electrochemical equipment consisted of a PGSTAT 12 or a μ-AUTOLAB (type III) driven by the GPES software. The working electrode consisted of a vitreous carbon disk that was polished on a felt tissue with alumina, rinsed with water, and dried before each CV scan. All of the potentials (text and figures) are referenced to the ferrocene−ferrocenium couple; ferrocene was added as an internal standard at the end of the experiments. The acetonitrile complex could not be obtained as a stable powder because of its decomposition in the absence of acetonitrile. For this complex, the electrochemical experiments were thus made by dilution (if needed) of solutions prepared just before the experiment. Although the amount of starting material (complex 1) and the volume of acetonitrile were controlled as precisely as possible, the concentration of 2 in the electrochemical experiments might not be as accurate as is usual. 4.2.1. Synthesis of [FeMo(CO)4(MeCN)(κ2-dppe)(μ-pdt)] (2). An acetonitrile solution (50 mL) of 1 (37 mg, 0.046 mmol) was heated under reflux. The progress of the reaction was followed by monitoring the decreasing intensity of the infrared carbonyl peaks of 1 and the concomitant increasing of the related peaks of the new product 2. After ca. 45 min, when the reaction was about to end, the solution turned from yellow-green to dark-green. Attempts to isolate the acetonitrile adduct 2 under as a powder failed because it decomposed rapidly in the absence of acetonitrile. In order to characterize 2 by NMR spectroscopy, this complex was also prepared in situ in a NMR tube by heating, at 100 °C, a CD3CN solution of 1 for 1 h 45 min. In other respects, crystals of [FeMo(CO)4(MeCN)(κ2-dppe)(μ-pdt)]. 2CH3CN (2·2CH3CN), suitable for X-ray analysis, were grown from a mixture of acetonitrile/diethyl ether in a 1:2 ratio, at −20 °C. L


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bridge Crystallographic Data Centre, 12 Union Road, Cambridge CB2 1EZ, UK; fax: +44 1223 336033.

Table 7. Crystallographic Data for Complex [FeMo(CO)4(CH3CN)(κ2-dppe)(μ-pdt)]·2CH3CN (2· 2CH3CN) empirical formula formula weight temp (K) cryst syst space group A (Å) B (Å) C (Å) β (deg) V (Å3) Z ρcalc (mg m−3) μ (mm−1) cryst size (mm) range of θ (deg) reflections measured Rint unique data/parameters R1 [I > 2σ(I)] R1 (all data) wR2 (all data) goodness of fit on F2 Δρmax, Δρmin/e Å−3

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C39H39FeMoN3O4P2S2 891.58 170(2) monoclinic P21/n 16.4001(6) 11.9850(3) 20.8457(6) 101.942(3) 4008.7(2) 4 1.477 0.902 0.50 × 0.10 × 0.05 2.89−26.37 30 030 0.0403 8178/477 0.0463 0.0615 0.1286 1.120 0.844, −0.442

*E-mail: [email protected] (M.B.). *E-mail: [email protected] (P.S.). *E-mail: [email protected] (J.T.). ORCID

Maurizio Bruschi: 0000-0002-5709-818X Philippe Schollhammer: 0000-0001-8161-7878 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS Financial support was provided by CNRS (Centre National de la Recherche Scientifique), Université de Bretagne Occidentale (UBO), and University of Milano-Bicocca. J.T. is grateful to Dr. M. Rudolph (University of Jena) for providing the DigiElch software. S.B. thanks UBO for providing a studentship. We are grateful to the X-ray and NMR departments of UBO for crystallographic measurements (Dr F. Michaud) and recording of variable temperature NMR experiments.

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REFERENCES

(1) Nicolet, Y.; Fontecilla-Camps, J. C. Fe−S Clusters : Biogenesis and Redox, Catalytic, and Regulatory Properties In Bioinspired Catalysis: Metal−Sulfur Complexes; Weigand, W., Schollhammer, P., Eds.; Wiley-VCH: Weinheim, 2015; Chapter 2, pp 21−48. (2) Rees, D. C. Great Metalloclusters in Enzymology. Annu. Rev. Biochem. 2002, 71, 221−246. (3) Johnson, D. C.; Dean, D. R.; Smith, A. D.; Johnson, M. K. Structure, function, and formation of biological iron-sulfur clusters. Annu. Rev. Biochem. 2005, 74, 247−281. (4) Fontecilla-Camps, J. C.; Volbeda, A.; Cavazza, C.; Nicolet, Y. Structure/Function Relationships of [NiFe]- and [FeFe]-Hydrogenases. Chem. Rev. 2007, 107, 4273−4303. (5) De Lacey, A. L.; Fernández, V. M.; Rousset, M.; Cammack, R. Activation and inactivation of hydrogenase function and the catalytic cycle: spectroelectrochemical studies. Chem. Rev. 2007, 107, 4304− 4330. (6) Lubitz, W.; Reijerse, E.; van Gastel, M. [NiFe] and [FeFe] Hydrogenases Studied by Advanced Magnetic Resonance Techniques. Chem. Rev. 2007, 107, 4331−4365. (7) Hu, Y.; Ribbe, M. W. Nitrogenase and homologs. J. Biol. Inorg Chem. 2014, 20, 435−445. (8) Burgess, B. K.; Lowe, D. J. Mechanism of Molybdenum Nitrogenase. Chem. Rev. 1996, 96, 2983−3012. (9) Thorneley, R. N. F.; Lowe, D. J. Nitrogenase: substrate binding and activation. J. Biol. Inorg Chem. 1996, 1, 576−580. (10) Spatzal, T.; Aksoyoglu, M.; Zhang, L.; Andrade, S. L. A.; Schleicher, E.; Weber, S.; Rees, D. C.; Einsle, O. Evidence for interstitial carbon in nitrogenase FeMo cofactor. Science 2011, 334, 940. (11) Lancaster, K. M.; Roemelt, M.; Ettenhuber, P.; Hu, Y.; Ribbe, M. W.; Neese, F.; Bergmann, U.; DeBeer, S. X-ray emission spectroscopy evidences a central carbon in the nitrogenase ironmolybdenum cofactor. Science 2011, 334, 974−977. (12) Hu, Y.; Ribbe, M. W. A journey into the active center of nitrogenase. J. Biol. Inorg Chem. 2014, 19, 731−736. (13) Einsle, O. Nitrogenase FeMo cofactor: an atomic structure in three simple steps. J. Biol. Inorg Chem. 2014, 19, 737−745. (14) Igarashi, R. Y.; Seefeldt, L. C. Nitrogen fixation: the mechanism of the Mo-dependent nitrogenase. Crit. Rev. Biochem. Mol. Biol. 2003, 38, 351−384.

for the SCF wavenumbers equal to 298.15 K, 1 bar, and 0.9914, respectively. Rotations have been treated classically, and vibrational modes have been described according to the harmonic approximation considering the frequency analysis carried out in vacuum. The effect of the solvent (dichloromethane, ε = 8.930; acetonitrile, ε = 36.64) has been evaluated according to the COSMO approach.97,98 Calculations of reduction potentials have been carried out using the Nernst equation, ΔG = −nFE, where n is the number of electron transferred and F is the Faraday constant. The resulting E value is an absolute reduction potential (Eabs), that is, it is not referenced to any standard electrode. Therefore, in order to obtain relative reduction potentials (Ecalc), the absolute reduction potential for the Fc+/Fc electrode, taken from ref 99 as 4.927 and 4.988 V for the CH2Cl2 and CH3CN solvents, respectively, has been systematically subtracted from Eabs values. Equilibrium constants have been calculated by first computing the ΔG of the reaction and then using the equation Kc = exp(−ΔG/RT) with T = 298.15 K. Both equilibrium constants and Ecalc values have been computed taking into account solvation contributions.

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ASSOCIATED CONTENT

* Supporting Information S

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.8b02861. Electrochemical data, cyclic voltammograms in the absence and in the presence of acids, scan rate dependence of current functions, digital simulations of cyclic voltammetric curves, rate constant of MeCN loss from 2−, and CIF file (PDF) Accession Codes

CCDC 1872311 contains the supplementary crystallographic data for this paper. These data can be obtained free of charge via www.ccdc.cam.ac.uk/data_request/cif, or by emailing [email protected], or by contacting The CamM


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Inorganic Chemistry

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