Electrochemistry of Simple Organometallic Models of Iron–Iron

Forum Article pubs.acs.org/IC

Electrochemistry of Simple Organometallic Models of Iron−Iron Hydrogenases in Organic Solvent and Water Frederic Gloaguen* UMR 6521, CNRS, Université de Bretagne Occidentale, CS 93837, 29238 Brest, France ABSTRACT: Synthetic models of the active site of iron−iron hydrogenases are currently the subjects of numerous studies aimed at developing H2-production catalysts based on cheap and abundant materials. In this context, the present report offers an electrochemist’s view of the catalysis of proton reduction by simple binuclear iron(I) thiolate complexes. Although these complexes probably do not follow a biocatalytic pathway, we analyze and discuss the interplay between the reduction potential and basicity and how these antagonist properties impact the mechanisms of proton-coupled electron transfer to the metal centers. This question is central to any consideration of the activity at the molecular level of hydrogenases and related enzymes. In a second part, special attention is paid to iron thiolate complexes holding rigid and unsaturated bridging ligands. The complexes that enjoy mild reduction potentials and stabilized reduced forms are promising iron-based catalysts for the photodriven evolution of H2 in organic solvents and, more importantly, in water.

■

INTRODUCTION The production of hydrogen (H2) from electrochemical reduction of a proton source, ideally water, represents an attractive approach for storing the electrical energy transiently produced by renewable energy sources such as sun and wind.1−3 Even though the hydrogen evolution reaction (HER) is efficiently catalyzed in aqueous acid solutions by platinumgroup metals (PGMs),4 there is a general consensus that high cost and limited natural abundance are obstacles to the use of PGMs in large-scale photoelectrochemical production of H2. In an attempt to overcome this problem, researchers have over the past decade revisited catalysis of the HER by earth-abundant transition-metal complexes,5−8 guided by the idea of a catalytic activity tailored by chemical design.9 While challenges remain, in particular the use of water as both a proton source and solvent,10 recent advances in this area have led to efficient HER catalysts employing cobalt11,12 or nickel.13 Iron is by far the most abundant transition metal on earth. The development of HER catalysts based on this metal therefore represents a significant goal to increase the cost efficiency of H2 production from water and renewable energy sources.14 The complex Fe(TPP) (TPP = tetraphenylpophyrin) catalyzes the HER,15 but the catalytically active iron(0) intermediate is generated in an organic solvent at an electrode potential more negative than −1.9 V (unless otherwise noted, all potentials are reported versus that of the Fc+/Fc couple), limiting practical interest. Iron macrocycles, for which the potential of the putative Fe0 oxidation state is positively shifted by electron-withdrawing groups, work at potentials milder than −1.3 V.16 Besides, catalytic oxidation of H2 has been established for an iron complex bearing phosphine ligands, in which a pendant amine group functions as a proton relay.17 © XXXX American Chemical Society

On the other hand, enzymes are a continuous source of inspiration for the design of synthetic catalysts. Several microorganisms produce and uptake H2 at high rate and low overpotential using hydrogenases (H2ases). The active sites of these enzymes contain nickel and/or iron.18,19 Unfortunately, sensitivity to oxygen has hampered until now large-scale use of H2ases. Yet, the crystallographic characterization of NiFe and FeFe H2ases, as well as a better understanding of their functions provided by spectroscopy,19 electrochemistry,20 and theoretical calculations,21 has opened new routes for developing synthetic HER catalysts based on iron. The biomimetic approach has been further facilitated by the finding that the active site of FeFe H2ases shares structural features with classical iron thiolate complexes of the type Fe2(SR)2(CO)6.22−24 As a result, numerous structural and functional models of FeFe H2ases have been synthesized and characterized.25−28 Recent developments in this area have led to iron thiolate complexes presenting salient features of the natural enzyme, such as a dithiolate bridge with a pendant amine group (adtH; Scheme 1), a terminal hydride, a rotated geometry, and a mixed FeIIFeI state.26,29 Catalytic oxidation of H2 has also been established for a synthetic model of the oxidized state of FeFe H2ase.30 Furthermore, an active artificial enzyme has been prepared by the incorporation of an iron thiolate complex into a scaffold protein,31 constituting the first step toward catalytic Special Issue: Small Molecule Activation: From Biological Principles to Energy Applications Part 3 Received: October 1, 2015

A

DOI: 10.1021/acs.inorgchem.5b02245 Inorg. Chem. XXXX, XXX, XXX−XXX

Forum Article

Inorganic Chemistry Scheme 1. Some of the Ligands Used in Iron Thiolate Models of FeFe H2ases

Figure 1. CVs at 0.5 V s−1 of 0.5 mM Fe2(bdt)(CO)5(P(OMe)3) in Bu4NPF6/MeCN before (blue) and after (red) the addition of 6 mM HOTs.

occur and to measure the catalytic current. A rigorous analysis of the CV responses under varied experimental conditions is, nevertheless, required to obtain key metrics of the catalytic activity, such as the overpotential and turnover frequency (TOF),48 and to determine the catalytic mechanism.49,50 Apart from electrochemical and coupled spectroscopic methods,51,52 computational chemistry methods are more and more employed to help analyze experimental data and decipher catalytic mechanisms.53 Density functional theory calculations have proved particularly useful to establish the structures of intermediates and transition states and for determining their thermodynamic properties, such as the acidity constant and standard electrochemical potential.54,55 Recently, computational chemistry methods have been coupled with kinetic analysis to calculate the TOF of catalysts in electrochemical reactions.56−58 Overpotential for Catalysis. A “good” HER catalyst exhibits a large catalytic current at a potential Ecat mildly more negative than the thermodynamic limit of the electrochemical production of H2 from a proton source HA in a given solvent, which is characterized by the standard potential E°HA/H2. The difference between these two potentials commonly defines the overpotential for catalysis. However, as recently discussed by Dempsey and co-workers,59 the significance of this parameter for quantitative evaluation depends on the method used to extract the value of Ecat from experimental data. It is advisible to derive Ecat from either the potential at which half of the catalytic peak current is obtained or the potential at which, in the absence of acid, the transition-metal complex reaches a catalytically relevant oxidation state. However, a more kinetically meaningful definition of the overpotential is given in eq 1, where E is the electrode potential required to drive the HER at a specific rate (i.e., at a given current).

activity improvements through chemical synthesis and mutagenesis. Several recent reviews have discussed in depth the chemistry of synthetic models of FeFe H2ases.29,32−35 The present report rather offers an electrochemist’s view of HER catalysis by simple binuclear iron(I) thiolate complexes that protonate at the metal−metal bond (Scheme 2) and as such probably do not follow a biocatalytic pathway, although the matter is debated.36,37 Scheme 2. Catalysis of the HER by a Simple FeFe H2ase Model That Protonates at the Metal−Metal Bond38

Herein, we first analyze and discuss the interplay between the reduction potential and basicity and how these antagonist properties impact the mechanisms of proton-coupled electron transfer (PCET) to the metal centers. This question is central to any consideration of the activity at the molecular level of H2ases and related enzymes.39−41 In a second part, special attention is paid to iron thiolate complexes holding rigid and unsaturated bridging ligands.42−44 The stability and mild reduction potential enjoyed by these types of complexes make them promising iron-based catalysts for the photodriven evolution of H2 in organic solvents and, more importantly, in water.45−47

■

RESULTS AND DISCUSSION Benchmarking the Performances of Homogeneous HER Catalysts. While water is the preferred proton source for the HER, many homogeneous catalysts are studied in organic solvents. This choice, often guided by solubility and stability issues, has several implications that are discussed below. Catalysis of the HER by transition-metal complexes is commonly evaluated by cyclic voltammetry (CV), where a catalytically inert electrode made of glassy carbon (GC) or mercury supplies electrons and an acid added to the organic solvent supplies protons. The catalytic activity is then indicated by changes of the CV responses with increasing acid concentrations. The main criterion is an enhancement of the reduction current at a potential close to that of the reduction of the complex (Figure 1) or its protonated form.38 CV can thus be employed to establish the potential at which catalysis does

η = E − E°HA/H2

(1)

The determination of E°HA/H2 has also been extensively discussed.60,61 This potential value is usually calculated from tabulated data using eq 2,62 in which E°H+/H2 is the standard potential of the solvated proton/hydrogen couple in a given solvent and pKa,HA the acidity constant of the proton source HA in the same solvent. E°HA/H2 = E°H+ /H2 − 2.3RT /F × pK a,HA

(2)

Recently, Bullock and co-workers reported that E°HA/H2 can be directly measured at an activated platinum electrode immersed in a H2-saturated solvent containing the proton source HA.63 Apparent Catalytic Rate Constant and TOFmax. It is of common practice to describe the catalysis of the HER by an B

DOI: 10.1021/acs.inorgchem.5b02245 Inorg. Chem. XXXX, XXX, XXX−XXX

Forum Article

Inorganic Chemistry ECcat mechanism, in which the fast electrochemical step E is followed by a rate-determining chemical step Ccat.13 The latter is characterized by an apparent rate constant kcat. Delahay and Stiehl64 developed the theory of ECcat mechanisms for polarography. Nicholson65 and Savéant66 independently developed a similar theory for CV. Under pseudo-first-order conditions, the catalytic current Icat is related to TOFmax = kcat[HA] (s−1) by the mathematical relationship shown in eq 3, where Ip is the peak current for reduction of the catalyst in the absence of acid HA and v is the potential scan rate used to record Icat and Ip (Figure 1). Icat RT × TOFmax = 4.48 Ip Fv

(3)

In principle, calculating TOFmax from eq 3 is restricted to situations where CV gives a catalytic wave that is S-shaped. However, most of the HER catalysts give peak-shaped responses displaying current hysteresis between the forward and return scans (Figure 1). This “nonideal” catalytic wave arises from depletion of acid, deactivation of catalyst, and/or inhibition by reaction products in the course of the CV scan. All of these side phenomena can interfere, lowering the measured catalytic current. Consequently, Savéant and co-workers67 recently proposed to extract TOFmax from the foot of the catalytic peak using a mathematical relationship that describes the electrochemical response under “pure” catalytic conditions (eq 4). At the beginning of the catalytic peak, the currents are low and therefore less altered by the side phenomena described above. 2.24 RT × TOFmax /Fv I = Ip 1 + exp[F(E − Ecat)/RT ]

Figure 2. Left: Background-corrected catalytic responses of Fe2(bdt)(CO)6 (top) and Fe2(bdt)(CO)5(P(OMe)3) (bottom) derived from CVs recorded at 0.5 V s−1 in the presence of 12 mol equiv of HOTs. Right: Foot-of-the-wave analyses. Linear fits (dotted line) to the data show deviation from “pure” catalytic conditions at I/Ip > 1.3.

(4) Figure 3. Plots of log TOF as a function of η = E − EHOTs/H2 for Fe2(bdt)(CO)6 (red) and Fe2(bdt)(CO)5(P(OMe)3) (blue) derived from the foot-of-the-wave analyses shown in Figure 2.

Note that eq 4 corresponds to limiting cases where the first chemical step is the rate-determining step.68 Case Study. To illustrate the approach described above, we compare Fe2(bdt)(CO)6 (Ecat = −1.28 V) with Fe2(bdt)(CO)5(P(OMe)3) (Ecat = −1.51 V) for catalysis of the reduction of tosic acid (HOTs; pKaMeCN = 8.7; E°HOTs/H2 = −0.65 V). Foot-of-the-wave analysis of the CV response of these two catalysts recorded in the presence of 12 mol equiv of HOTs is shown in Figure 2. Deviation from linearity (i.e., “pure” catalytic condition) occurs at I/Ip > 1.3. From the slope of the linear portion of the plot, eq 4 gives TOFmax = 41 and 2400 s−1 for the all-CO and P(OMe)3 derivatives, respectively. Now, the question that arises is whether the larger value of TOFmax calculated for the P(OMe)3 derivative is due to an enhanced basicity of the iron−iron site accelerating the protonation steps or to a more negative value of Ecat increasing the driving force (i.e., Ecat − E°HOTs/H2) for the electrochemical reaction. As proposed by Savéant and co-workers,67 this issue can be addressed by establishing a Tafel-like relationship between log TOF and η (eq 5).

TOF0 and is intrinsically less efficient than the all-CO derivative for HER catalysis from HOTs. Although not yet experimentally verified, we argue that an opposite result might have been obtained with a weaker acid as a proton source, i.e., with an acid for which the E°H2/HA value is closer to the Ecat value of the P(OMe)3 derivative. Catalysis of the HER by Transition-Metal Complexes. Metal hydride complexes have long attracted interest because of their possible role in H+/H2 activation through proton/hydride coupling, i.e., H+ + H− → H2.69 Cycles based on this concept can also describe catalysis of the HER by iron thiolate complexes.70 The key intermediate is the metal hydride M−H (Scheme 3). Depending on the basicity of the metal center in Scheme 3. Catalytic Cycles for the Reduction of Protons to H2 by a Transition-Metal Complex M

log TOF = log TOFmax + F(Ecat − E°HA/ H 2 )/2.3RT − Fη /2.3RT

(5)

Extrapolating plots of log TOF versus η give the intrinsic efficiency of the catalytic systems at zero overpotential (Figure 3): i.e., TOF0 = 9.3 × 10−8 and 7.0 × 10−10 M−1 s−1 for the allCO and P(OMe)3 derivatives, respectively. Hence, despite achieving a larger TOFmax, the P(OMe)3 derivative has a lower C

DOI: 10.1021/acs.inorgchem.5b02245 Inorg. Chem. XXXX, XXX, XXX−XXX

Forum Article

Inorganic Chemistry

that the hydride intermediate is then engaged in a bimolecular reaction that releases H2, regenerating the starting complex. Stepwise versus Concerted PCET. As discussed above, catalysis of the HER by transition-metal complexes is thermodynamically and kinetically controlled by the reactivity of a metal hydride intermediate. This species can be formed upon PCET to the metal center, which means here that electron transfer results in an increase of the basicity of the metal center, allowing its protonation. At least one additional PCET step is then required to release H2. Electron and proton transfer are frequently supposed to proceed in a stepwise fashion (Scheme 3). However, one can also consider a concerted proton−electron transfer (CPET), avoiding the formation of highly reduced species or the use of strong acid as the proton source and thus enhancing the catalysis.74 CPET mechanisms are likely to occur in catalysis of the HER by FeFe H2ase models, especially those holding ligands with a pendant base such as adtH, but this has yet to be established. In collaboration with Hammarström and co-workers,75 we have recently reported the first experimental evidence that oxidative cleavage of a metal−hydride bond can follow a CPET mechanism at a low overall driving force (Scheme 5). No

M and on the strength of the acid being used as a proton source, M−H can be either a stable complex formed by the oxidative addition of a proton (step b) or a transient species generated via PCET (steps a and b). The possibility of different pathways is also to be considered in the subsequent catalytic steps that are not shown in Scheme 3. Rerouting from one pathway to another is dependent upon a combination of thermodynamic and kinetic factors that are discussed below with selected examples. Closing the Catalytic Cycle. In catalysis of the HER by transition-metal complexes, an important point to consider is the possibility of closing the cycle by facile release of H2 from metal hydrides. As shown in Scheme 3, reduction and protonation of M−H give a two-electron two-proton intermediate M−2H (step c), which can be either a dihydrogen complex M−(η2-H2) or a dihydride M−(H)2. There are numerous examples showing that for mononuclear transition-metal complexes the kinetic site of protonation in M−H− is the hydride ligand not the metal center.71 However, instead of releasing H2 (step d), the M−(η2-H2) intermediate can evolve toward a stable dihydride, from which subsequent formation and elimination of H2 is more demanding. Similarly, it has been proposed that reduction and subsequent protonation of the bridging hydride [(H)Fe2(pdt)(CO)4(dppe)]+ gives a kinetically favored η2-H2 intermediate that transforms into a stable dihydride on the CV time scale (Scheme 4).72 Dihydrogen complexes can be

Scheme 5. Stepwise (Black) versus Concerted (Red) Pathways for Oxidative Cleavage of the Metal−Hydride Bond in (H)W(Cp)(CO)3 (M−H)

Scheme 4. Proposed Mechanisms for H2 Release from an Iron Thiolate Hydride Intermediate, in Which the Hydride Ligand Is in the Bridging Position

catalysis is involved here, but demonstration of a CPET mechanism in the reactivity of the metal hydride complex presents new opportunities for the development of efficient HER catalysts. Interplay between the Reduction Potential and Basicity. The increase of the basicity of the binuclear iron(I) site upon the substitution of electron-donating ligands for carbonyls in iron thiolate complexes has been employed by Poilblanc and co-workers to prepare stable bridging hydrides [(H)Fe2(SR)2(CO)4(PR3)2]+ (R = alkyl).76 Rauchfuss and coworkers have further shown that reduction of the structurally related hydride [(H)Fe2(pdt)(CO)4(PMe3)2]+ catalyzes the HER in MeCN when an excess of a relatively strong acid (pKaMeCN < 10) is added to the solution (Scheme 2).38,77 Chemical and electrochemical data support a catalytic mechanism that begins with the oxidative addition of a proton to the binuclear iron(I) site to give a bridging hydride, the reduction of which at Ep,red = −1.37 V gives a product able to react with a second proton. Whether the release of H2 precedes or is followed by the second electron transfer to the complex remains to be established. Importantly, the negative shift of the reduction potentials resulting from the replacement of two CO ligands by two electron-donor ligands, such as PMe3, is compensated for by a positive shift of about 1 V of the reduction potential of the corresponding iron thiolate hydride, consistent with an oxidative addition of a proton to the iron−iron site.38,77 From the thermodynamic cycle shown in Scheme 3, it can be seen that the pKa values and the reversible potentials E° of the

acidic, which suggests that cleavage of the hydrogen−hydrogen bond may arise here from proton transfer from the H2 ligand to the “electron-rich” iron centers. Formation and release of H2 from the dihydride require then an additional electron and/or proton transfer, which has a significant thermodynamic cost. Interestingly, the formation of a stable dihydride intermediate seems to be less favored when the iron thiolate complex holds a terminal hydride ligand and an adtH bridge (see the Biocatalytic Pathway section).29 An alternative pathway for the release of H2 is the bimolecular combination of two M−H intermediates (step e, Scheme 3). This pathway has been proposed to explain the catalysis of H2 by Fe2(pdt)(CO)5(P(OMe)3) at a potential of about 300 mV less negative than that of reduction of the complex in the absence of acid.73 CV in the presence of HOTs shows that the hydride intermediate (H)Fe2(pdt)(CO)5(P(OMe)3) is formed via a CE mechanism in the vicinity of the electrode. Controlled potential electrolysis experiments suggest D

DOI: 10.1021/acs.inorgchem.5b02245 Inorg. Chem. XXXX, XXX, XXX−XXX

Forum Article

Inorganic Chemistry Scheme 6. Catalysis of the HER by Fe2(bdt)(CO)6 from Either a Strong or a Weak Acid as the Proton Source

protonated species (i.e., M−H and M−H−) and unprotonated species (i.e., M− and M2−) are interdependent (eq 5).

iron axis favoring the formation of a reactive terminal hydride. On that basis, Rauchfuss and co-workers synthesized Fe2(adtH)(CO)2(dppv)2 (adtH and dppv; Scheme 1),78 a FeFe H2ase model that catalyzes the HER in organic solvent at an unprecedented TOF of 58000 s−1 with an overpotential of −0.51 V. The advantages of a pendant base and a rotated symmetry to accelerate the HER have been extensively discussed in recent reviews,29,33 to which the reader is referred for further details. Note, however, that the catalytic activity of Fe2(adtH)(CO)2(dppv)2 is significantly altered in the presence of coordinating solvent and ions, indicating that precise control of the chemical species approaching the iron thiolate core will be necessary for practical application of catalysts working via a biocatalytic pathway. This role is devoted to the scaffold protein in natural H2ases.79 Catalysis of the HER by All-CO Iron Thiolate Complexes. Darensbourg and co-workers first reported catalysis of the HER from a weak acid such as acetic acid (HOAc; pKaMeCN = 22.3, E°HOAc/H2 = −1.46 V) by the initial reduction of Fe2(pdt)(CO)6 at Ep,red = −1.67 V.80 An EECC mechanism has been initially proposed in which the reduction of the all-CO iron(I) thiolate complex to the formal Fe0Fe0 state significantly increases the basicity of the iron−iron site, allowing its reaction with a weak acid. This assumption is fully supported by the previous finding that reduction of an iron thiolate hydride from the FeIIFeII to FeIFeII state results in an increase of its basicity by ca. 17 pKa units (see the Interplay between the Reduction Potential and Basicity section). Pickett and co-workers81 thoroughly discussed the question of the number of electrons involved in the primary reduction of all-CO iron thiolate complexes in the absence of acid. Electrochemical, spectroelectrochemical, and theoretical studies have thus revealed that the electronic and steric properties of the thiolate bridge have a subtle influence on the electrochemical behavior, emphasizing the variety of the products that result from initial electron transfer to all-CO iron thiolate complexes.26,70 All-CO Complexes with Rigid and Conjugated Thiolate as the Bridging Ligand. Obviously, all-CO iron thiolate complexes giving stable Fe0FeI or Fe0Fe0 forms will be more inclined to achieve a clean HER catalysis without side reactions. To this aim, it has been thought to replace flexible alkyl thiolates (e.g., pdt or adtH; Scheme 1) with rigid and conjugated thiolates (e.g., bdt, dtCO2Me, or bpdt; Scheme

E°(M−H/M−H−) − E°(M−/M2 −) = 2.3RT /F[pK a(M−H−) − pK a(M−H)]

(5)

The positive shift of 1 V between the reduction potentials of [(H)Fe2(pdt)(CO)4(PMe3)2]+ and Fe2(pdt)(CO)4(PMe3)2 allows one thus to estimate that electron transfer results in an increase of the pKa value of iron thiolate hydride by ca. 17 units. In order to protonate at the iron−iron bond, an iron thiolate complex must be sufficiently “electron-rich”, which is achieved by substitution of electron-donating ligands (e.g., phosphine) for carbonyls in the iron coordination sphere. In the meantime, the complex becomes more difficult to reduce, and the positive potential shift resulting from the formation of iron thiolate hydride is partly or totally annihilated by the negative potential shift resulting from the introduction of strong electrondonating ligands.70 One possibility to reconcile the above contradictory conditions is to introduce a protanable site, which would have little effect on the reduction potential of the iron thiolate complex. This might happen when the lowest unoccupied molecular orbital (LUMO) of the complex receives little contribution from the protonable site, which seems to be the case when an amino group is introduced in the dithiolate ligand (see the Biocatalytic Pathway section). Another possibility is to increase the basicity of the iron−iron site by electron transfer to an easily reduced, and hence nonprotonable, iron thiolate complex, in which the LUMO still has significant Fe−Fe bond character (see the Catalysis of the HER by All-CO Iron Thiolate Complexes section). Biocatalytic Pathway. Substituted iron thiolate complexes of the type Fe2(SR)2(CO)6−xLx (L = electron-donating ligand and x ≤ 4) are clearly relevant to the active site of FeFe H2ases, where cyanide ligands help to stabilize the low oxidation states of the iron−iron center. However, analysis of the CV data indicates that the rate-determining step is the formation of the bridging hydride (Scheme 2).38 This result is not surprising considering the kinetic barriers to protonation of the metal centers. To overcome this problem, two biomimetic features have been identified and targeted. First, a protonable amino group in the dithiolate ligand shuttling the proton to the “electron-rich” metal centers and second a rotated geometry along the iron− E

DOI: 10.1021/acs.inorgchem.5b02245 Inorg. Chem. XXXX, XXX, XXX−XXX

Forum Article

Inorganic Chemistry 1),49,82−85 for which a greater ligand−metal mixing of the frontier orbitals is anticipated.86 In acetonitrile, Fe2(bdt)(CO)6 is reduced at E1/2 = −1.28 V to its dianion in a chemically reversible two-electron process (E2 > E1; Scheme 6) because of the large structural rearrangement that follows the first electron transfer at E1.49,83 In comparison, Fe2(pdt)(CO)6, which holds a flexible alkyl dithiolate as the bridging ligand, undergoes a partially reversible reduction at a potential of ca. 0.4 V more negative. Thus, the bdt ligand is able to both modulate the reduction potentials and stabilize the reduced forms of the iron−iron site. In the presence of moderately strong acids such as HOTs, the reduction of Fe2(bdt)(CO)6 becomes chemically irreversible, indicating a reaction of the reduced form with acid. There is a further increase of the reduction peak current with increasing acid concentrations, confirming that catalytic reduction of HOTs occurs at this potential.49 The complex Fe2(bdt)(CO)6 is also able to catalyze the reduction of a weak acid such as HOAc83 but at a significantly more negative potential of approximately −2 V (E3; Scheme 6). Contrary to HOTs, HOAc is too weak to protonate the twoelectron one-proton intermediate. The latter must be first reduced around −2 V before protonation can occur. The ensuing fast release of H2 regenerates the one-electron-reduced intermediate, which is the catalytically active species for the reduction of a weak acid. Contrary to the bdt ligand, several dithiolate ligands featuring electron-withdrawing groups such as o-carborane,84 naphthalene,87 or biphenyl88,89 favor the transfer of two electrons in two separated one-electron steps. The all-CO binuclear iron(I) complexes bearing a benzo[c]cinnoline give also a similar electrochemical response.90 It is usually postulated that the first reduction step gives a stabilized FeIFe0 intermediate. However, a significant contribution of the bridging ligand to the LUMO cannot be completely ruled out here. For example, the two reduction steps of the dithiolene derivative Fe2(dtCO2Me)(CO)6 at E1/2 = −1.11 and −1.25 V do not trigger the catalysis of proton reduction from a strong acid (HBF4/Et2O, pKaMeCN ∼ 3, and E°HBF4/H2 ∼ −0.32 V) despite the large driving force for this reaction.89 This observation suggests that the reduction of Fe2(dtCO2Me)(CO)6 does not result in a large increase of the basicity of the iron−iron site because the LUMO of this complex does not have significant iron−iron bond character. Similarly, the primary reduction of Fe2(pdt)(CO)4(bma), in which bma is a redox-active ligand chelating one iron center (Scheme 1), occurs at a potential 0.74 V milder than that of the all-CO parent compound.91 This result is consistent with a first electron transfer to the bma ligand. However, CV in the presence of a strong acid shows no signs of proton reduction catalysis. These observations strongly suggest a weak electronic communication between the bma ligand and the iron−iron site, preventing an increase of its basicity upon electron transfer and hence the formation of a metal hydride intermediate. Electrochemical and Photodriven HER Catalysis by Iron Thiolate Complexes in Water. The lack of solubility of the majority of the FeFe H2ase models in aqueous media has limited their study to organic solvents. Nevertheless, a few iron thiolate complexes holding hydrophilic ligands have been prepared, confirming stability and activity of these types of catalysts in mixtures of organic solvents and water.92,93 Darensbourg and co-workers, on the other hand, showed that a sulfonated iron thiolate derivative could be included in the

cavity of a water-soluble cyclodextrin.94 This approach solves the problem of water solubility and further offers a possibility of mimicking the protein environment of H2ases.95,96 Unfortunately, the iron thiolate/cyclodextrin system does not catalyze the HER in water, when electrons are supplied by a GC electrode. In parallel, few studies concerning photodriven H 2 production by iron thiolate catalysts in aqueous solution containing a fraction of organic solvent have been reported.97−99 Interestingly, Wu and co-workers100 have shown that water-insoluble iron thiolate complexes could be incorporated into micelles formed in an aqueous sodium dodecylsulfate (SDS) solution. This system produces H2 with low rates using a PGM-based photosensitizer and ascorbic acid as a sacrificial electron donor. However, the catalytic properties of the iron thiolate complexes have not been independently evaluated in aqueous SDS solution. Whether the H2 production rate is limited by the activity of the catalyst, the coupling with the photosensitizer, or any other process has therefore not been established. These results prompted us to study the electrochemistry of the water-insoluble complex Fe2(bdt)(CO)6 in aqueous SDS solution.45 The bdt ligand is expected to decrease the reduction potential of the complex and to stabilize the catalytically active forms Fe0FeI and Fe0Fe0 (see above). Polarography was employed to ensure fast electron transfer between the electrode and Fe2(bdt)(CO)6 included in the SDS micelles. At neutral pH, the primary reduction of the complex occurs at E1/2 = −0.74 V vs SHE, a potential value of about 0.6 V less negative than those reported for iron thiolate complexes bearing hydrophilic ligands.92,93 The linear dependence of the reduction peak current on the concentration of Fe2(bdt)(CO)6 in the range 20−60 μM, moreover, indicates the reduction of a free-diffusing species, confirming that there is no adsorption of the complex on the mercury electrode. The addition of 10 mM HOAc (pH 3.3) triggers a large current increase at a potential slightly less negative than that of the primary reduction of Fe2(bdt)(CO)6 at neutral pH. Because, at pH 3.3, the direct reduction of free protons on mercury occurs below −1.2 V vs SHE, the current increase around −0.7 V vs SHE is therefore ascribed to the catalytic reduction of HOAc mediated by Fe2(bdt)(CO)6. Analysis of polarographic waves under electrocatalytic conditions gives a value of TOFmax ∼ 2600 s−1 at an overpotential of 0.5 V. Encouraged by these results, we reasoned that Fe2(bdt)(CO)6 could be an efficient proton reduction catalyst for photodriven HER in water. We thus analyzed the performance of a PGM-free photocatalytic system consisting of Fe2(bdt)(CO)6 as a HER catalyst, Eosin Y (EY2−) as a photosensitizer, and triethylamine (Et3N) as a sacrificial electron donor dissolved in aqueous SDS solution at pH 10.5.47 The amount of H2 formed upon irradiation of this photocatalytic system with a light-emitting diode (λ = 455 nm) exceeds 0.8 mL in 4.5 h, corresponding to a TON of 117 mol of H2 per mol of Fe2(bdt)(CO)6. The lifetime of the photocatalytic system is clearly limited by the stability of the photosensitizer. The production of H2 lasts more than 30 h when the molecular ratio Eosin Y/catalyst is increased from 2 to 10. Besides, the initial rate of the photodriven HER (TOF ∼ 60 s−1) is significantly smaller than that measured under electrocatalytic conditions (TOF ∼ 2600 s−1). Thermodynamic considerations and UV/vis spectroscopy experiments suggest that the catalytic cycle begins with the F

DOI: 10.1021/acs.inorgchem.5b02245 Inorg. Chem. XXXX, XXX, XXX−XXX

Forum Article

Inorganic Chemistry

(2) Gray, H. B. Nat. Chem. 2009, 1, 7. (3) Thapper, A.; Styring, S. r.; Saracco, G.; Rutherford, A. W.; Robert, B.; Magnuson, A.; Lubitz, W.; Llobet, A.; Kurz, P.; Holzwarth, A.; Fiechter, S.; de Groot, H.; Campagna, S.; Braun, A.; Bercegol, H.; Artero, V. Green 2013, 3, 43. (4) Esposito, D. V.; Hunt, S. T.; Stottlemyer, A. L.; Dobson, K. D.; McCandless, B. E.; Birkmire, R. W.; Chen, J. G. Angew. Chem., Int. Ed. 2010, 49, 9859−9862. (5) Rakowski Dubois, M.; Dubois, D. L. Acc. Chem. Res. 2009, 42, 1974−1982. (6) Artero, V.; Chavarot-Kerlidou, M.; Fontecave, M. Angew. Chem., Int. Ed. 2011, 50, 7238−7266. (7) Wang, M.; Chen, L.; Sun, L. Energy Environ. Sci. 2012, 5, 6763− 6778. (8) Queyriaux, N.; Jane, R. T.; Massin, J.; Artero, V.; ChavarotKerlidou, M. Coord. Chem. Rev. 2015, 304−305, 3−19. (9) Crabtree, R. H. New J. Chem. 2011, 35, 18−23. (10) Thoi, V. S.; Karunadasa, H. I.; Surendranath, Y.; Long, J. R.; Chang, C. J. Energy Environ. Sci. 2012, 5, 7762−7770. (11) Sun, Y.; Bigi, J. P.; Piro, N. A.; Tang, M. L.; Long, J. R.; Chang, C. J. J. Am. Chem. Soc. 2011, 133, 9212−9215. (12) Chen, L.; Wang, M.; Han, K.; Zhang, P.; Gloaguen, F.; Sun, L. Energy Environ. Sci. 2014, 7, 329−334. (13) Helm, M. L.; Stewart, M. P.; Bullock, R. M.; DuBois, M. R.; DuBois, D. L. Science 2011, 333, 863−866. (14) Simmons, T. R.; Artero, V. Angew. Chem., Int. Ed. 2013, 52, 6143−6145. (15) Bhugun, I.; Lexa, D.; Savéant, J.-M. J. Am. Chem. Soc. 1996, 118, 3982−3983. (16) Rose, M. J.; Gray, H. B.; Winkler, J. R. J. Am. Chem. Soc. 2012, 134, 8310−8313. (17) Liu, T.; DuBois, D. L.; Bullock, R. M. Nat. Chem. 2013, 5, 228− 233. (18) Fontecilla-Camps, J. C.; Volbeda, A.; Cavazza, C.; Nicolet, Y. Chem. Rev. 2007, 107, 4273−4303. (19) Lubitz, W.; Ogata, H.; Rudiger, O.; Reijerse, E. Chem. Rev. 2014, 114, 4081−4148. (20) Hexter, S. V.; Grey, F.; Happe, T.; Climent, V.; Armstrong, F. A. Proc. Natl. Acad. Sci. U. S. A. 2012, 109, 11516−11521. (21) Fourmond, V.; Greco, C.; Sybirna, K.; Baffert, C.; Wang, P.-H.; Ezanno, P.; Montefiori, M.; Bruschi, M.; Meynial-Salles, I.; Soucaille, P.; Blumberger, J.; Bottin, H.; De Gioia, L.; Leger, C. Nat. Chem. 2014, 6, 336−342. (22) Lyon, E. J.; Georgakaki, I. P.; Reibenspies, J. H.; Darensbourg, M. Y. Angew. Chem., Int. Ed. 1999, 38, 3178−3180. (23) Le Cloirec, A.; Davies, S. C.; Evans, D. J.; Hughes, D. L.; Pickett, C. J.; Best, S. P.; Borg, S. Chem. Commun. 1999, 2285−2286. (24) Schmidt, M.; Contakes, S. M.; Rauchfuss, T. B. J. Am. Chem. Soc. 1999, 121, 9736−9737. (25) Tard, C.; Liu, X.; Ibrahim, S. K.; Bruschi, M.; Gioia, L. D.; Davies, S. C.; Yang, X.; Wang, L.-S.; Sawers, G.; Pickett, C. J. Nature 2005, 433, 610−613. (26) Capon, J. F.; Gloaguen, F.; Petillon, F. Y.; Schollhammer, P.; Talarmin, J. Coord. Chem. Rev. 2009, 253, 1476−1494. (27) 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. (28) Barton, B. E.; Olsen, M. T.; Rauchfuss, T. B. Curr. Opin. Biotechnol. 2010, 21, 292−297. (29) Rauchfuss, T. B. Acc. Chem. Res. 2015, 48, 2107−2116. (30) Camara, J. M.; Rauchfuss, T. B. Nat. Chem. 2011, 4, 26−30. (31) Berggren, G.; Adamska, A.; Lambertz, C.; Simmons, T. R.; Esselborn, J.; Atta, M.; Gambarelli, S.; Mouesca, J. M.; Reijerse, E.; Lubitz, W.; Happe, T.; Artero, V.; Fontecave, M. Nature 2013, 499, 66−69. (32) Ghosh, S.; Hogarth, G.; Hollingsworth, N.; Holt, K. B.; Richards, I.; Richmond, M. G.; Sanchez, B. E.; Unwin, D. Dalton Trans. 2013, 42, 6775−6792.

photodriven reduction of Fe2(bdt)(CO)6 (Scheme 7). The reduced Fe0FeI intermediate reacts with a proton source to Scheme 7. Proposed Mechanism for the Formation of an Iron Thiolate Hydride Intermediate under Photocatalytic Conditions

yield iron hydride. Subsequent reduction and protonation steps produce H2, regenerating the starting complex. More studies are, however, needed to identify the proton source and to gain further insight into the reaction mechanism.

■

CONCLUSIONS The simple iron thiolate complexes described herein catalyze the reduction of a proton source to H2 without following a biocatalytic pathway. However, they represent good models to study the interplay between the reduction potential and basicity and how these antagonist properties impact the mechanisms of PCET to metal centers, which is a central question to any consideration of the activity at the molecular level of H2ases and related enzymes. Besides, we have shown that analysis of the CV responses using rigorous kinetic models provides key metrics of the catalytic activity, such as the overpotential, TOFmax, and TOF0. This might help our understanding of the electronic and steric effects of the ligands on the reactivity of the iron−iron center toward protons. Finally, host structures that incorporate iron thiolate complexes offer exciting opportunities to build catalytic systems that produce H2 from water. However, difficult issues have to be addressed to rationalize the choice of the host structure. For example, determining the impact of the electrostatic, hydrophobic, and steric properties of the host structure on the reactivity of the iron thiolate complex will require basic knowledge of how PCET proceeds between different chemical phases and/or in a confined environment.

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS This research was supported by the CNRS, UBO, and ANR. We are indebted to the contributions of co-workers whose names are cited in the publications from our laboratory. Special thanks are due to Prof. Mei Wang (Dalian), Prof. Licheng Sun (Stockholm), and Prof. Thomas Rauchfuss (Urbana−Champaign) for continued cooperation.

■

REFERENCES

(1) Lewis, N. S.; Nocera, D. G. Proc. Natl. Acad. Sci. U. S. A. 2006, 103, 15729−15735. G

DOI: 10.1021/acs.inorgchem.5b02245 Inorg. Chem. XXXX, XXX, XXX−XXX

Forum Article

Inorganic Chemistry

(68) Costentin, C.; Savéant, J.-M. ChemElectroChem 2015, 2, 1774− 1784. (69) Koelle, U.; Paul, S. Inorg. Chem. 1986, 25, 2689−2694. (70) Capon, J. F.; Gloaguen, F.; Schollhammer, P.; Talarmin, J. Coord. Chem. Rev. 2005, 249, 1664−1676. (71) Gordon, J. C.; Kubas, G. J. Organometallics 2010, 29, 4682− 4701. (72) Ezzaher, S.; Capon, J. F.; Dumontet, N.; Gloaguen, F.; Petillon, F. Y.; Schollhammer, P.; Talarmin, J. J. Electroanal. Chem. 2009, 626, 161−170. (73) Morvan, D.; Capon, J. F.; Gloaguen, F.; Schollhammer, P.; Talarmin, J. Eur. J. Inorg. Chem. 2007, 2007, 5062−5068. (74) Costentin, C.; Robert, M.; Saveant, J.-M. Acc. Chem. Res. 2010, 43, 1019−1029. (75) Bourrez, M.; Steinmetz, R.; Ott, S.; Gloaguen, F.; Hammarstrom, L. Nat. Chem. 2015, 7, 140−145. (76) Fauvel, K.; Mathieu, R.; Poilblanc, R. Inorg. Chem. 1976, 15, 976−978. (77) Gloaguen, F.; Lawrence, J. D.; Rauchfuss, T. B. J. Am. Chem. Soc. 2001, 123, 9476−9477. (78) Carroll, M. E.; Barton, B. E.; Rauchfuss, T. B.; Carroll, P. J. J. Am. Chem. Soc. 2012, 134, 18843−18852. (79) Artero, V.; Berggren, G.; Atta, M.; Caserta, G.; Roy, S.; Pecqueur, L.; Fontecave, M. Acc. Chem. Res. 2015, 48, 2380−2387. (80) Chong, D.; Georgakaki, I. P.; Mejia-Rodriguez, R.; SanabriaChinchilla, J.; Soriaga, M. P.; Darensbourg, M. Y. Dalton Trans. 2003, 4158−4163. (81) Borg, S. J.; Behrsing, T.; Best, S. P.; Razavet, M.; Liu, X.; Pickett, C. J. J. Am. Chem. Soc. 2004, 126, 16988−16999. (82) Capon, J. F.; Gloaguen, F.; Schollhammer, P.; Talarmin, J. J. Electroanal. Chem. 2004, 566, 241−247. (83) Felton, G. A. N.; Vannucci, A. K.; Chen, J.; Lockett, L. T.; Okumura, N.; Petro, B. J.; Zakai, U. I.; Evans, D. H.; Glass, R. S.; Lichtenberger, D. L. J. Am. Chem. Soc. 2007, 129, 12521−12530. (84) Schwartz, L.; Eriksson, L.; Lomoth, R.; Teixidor, F.; Vinas, C.; Ott, S. Dalton Trans. 2008, 2379−2381. (85) Kaur-Ghumaan, S.; Schwartz, L.; Lomoth, R.; Stein, M.; Ott, S. Angew. Chem., Int. Ed. 2010, 49, 8033−8036. (86) Eisenberg, R.; Gray, H. B. Inorg. Chem. 2011, 50, 9741−9751. (87) Wright, R. J.; Lim, C.; Tilley, T. D. Chem. - Eur. J. 2009, 15, 8518−8525. (88) Singh, P. S.; Rudbeck, H. C.; Huang, P.; Ezzaher, S.; Eriksson, L.; Stein, M.; Ott, S.; Lomoth, R. Inorg. Chem. 2009, 48, 10883− 10885. (89) Charreteur, K.; Kdider, M.; Capon, J. F.; Gloaguen, F.; Petillon, F. Y.; Schollhammer, P.; Talarmin, J. Inorg. Chem. 2010, 49, 2496− 2501. (90) Orain, P. Y.; Capon, J. F.; Gloaguen, F.; Schollhammer, P.; Talarmin, J. Int. J. Hydrogen Energy 2010, 35, 10797−10802. (91) Si, Y. T.; Charreteur, K.; Capon, J. F.; Gloaguen, F.; Petillon, F. Y.; Schollhammer, P.; Talarmin, J. J. Inorg. Biochem. 2010, 104, 1038− 1042. (92) Mejia-Rodriguez, R.; Chong, D.; Reibenspies, J. H.; Soriaga, M. P.; Darensbourg, M. Y. J. Am. Chem. Soc. 2004, 126, 12004−12014. (93) Na, Y.; Wang, M.; Jin, K.; Zhang, R.; Sun, L. J. Organomet. Chem. 2006, 691, 5045−5051. (94) Singleton, M. L.; Reibenspies, J. H.; Darensbourg, M. Y. J. Am. Chem. Soc. 2010, 132, 8870−8871. (95) Ginovska-Pangovska, B.; Dutta, A.; Reback, M. L.; Linehan, J. C.; Shaw, W. J. Acc. Chem. Res. 2014, 47, 2621−2630. (96) Caserta, G.; Roy, S.; Atta, M.; Artero, V.; Fontecave, M. Curr. Opin. Chem. Biol. 2015, 25, 36−47. (97) Li, X.; Wang, M.; Chen, L.; Wang, X.; Dong, J.; Sun, L. ChemSusChem 2012, 5, 913−919. (98) Li, X.; Wang, M.; Zheng, D.; Han, K.; Dong, J.; Sun, L. Energy Environ. Sci. 2012, 5, 8220−8224. (99) Wang, F.; Wang, W.-G.; Wang, H.-Y.; Si, G.; Tung, C.-H.; Wu, L.-Z. ACS Catal. 2012, 2, 407−416.

(33) Wang, N.; Wang, M.; Chen, L.; Sun, L. Dalton Trans. 2013, 42, 12059−12071. (34) Simmons, T. R.; Berggren, G.; Bacchi, M.; Fontecave, M.; Artero, V. Coord. Chem. Rev. 2014, 270−271, 127−150. (35) Apfel, U.-P.; Pétillon, F. Y.; Schollhammer, P.; Talarmin, J.; Weigand, W. [FeFe] Hydrogenase Models: an Overview. In Bioinspired Catalysis; Schollhammer, P., Weigand, W., Eds.; WileyVCH Verlag: Berlin, 2014; pp 79−104. (36) Lindahl, P. A. J. Inorg. Biochem. 2012, 106, 172−178. (37) Chernev, P.; Lambertz, C.; Bruenje, A.; Leidel, N.; Sigfridsson, K. G. V.; Kositzki, R.; Hsieh, C.-H.; Yao, S.; Schiwon, R.; Driess, M.; Limberg, C.; Happe, T.; Haumann, M. Inorg. Chem. 2014, 53, 12164− 12177. (38) Gloaguen, F.; Lawrence, J. D.; Rauchfuss, T. B.; Benard, M.; Rohmer, M. M. Inorg. Chem. 2002, 41, 6573−6582. (39) Mulder, D. W.; Ratzloff, M. W.; Bruschi, M.; Greco, C.; Koonce, E.; Peters, J. W.; King, P. W. J. Am. Chem. Soc. 2014, 136, 15394− 15402. (40) Albers, A.; Demeshko, S.; Dechert, S.; Saouma, C. T.; Mayer, J. M.; Meyer, F. J. Am. Chem. Soc. 2014, 136, 3946−3954. (41) Dempsey, J. L. Nat. Chem. 2015, 7, 101−102. (42) Chen, L.; Wang, M.; Gloaguen, F.; Zheng, D.; Zhang, P.; Sun, L. Chem. - Eur. J. 2012, 18, 13968−13973. (43) Chen, L.; Wang, M.; Gloaguen, F.; Zheng, D.; Zhang, P.; Sun, L. Inorg. Chem. 2013, 52, 1798−1806. (44) Pandey, I. K.; Mobin, S. M.; Deibel, N.; Sarkar, B.; KaurGhumaan, S. Eur. J. Inorg. Chem. 2015, 2015, 2875−2882. (45) Quentel, F.; Passard, G.; Gloaguen, F. Energy Environ. Sci. 2012, 5, 7757−7761. (46) Quentel, F.; Passard, G.; Gloaguen, F. Chem. - Eur. J. 2012, 18, 13473−13479. (47) Orain, C.; Quentel, F.; Gloaguen, F. ChemSusChem 2014, 7, 638−643. (48) Quentel, F.; Gloaguen, F. Electrochim. Acta 2013, 110, 641−645. (49) Capon, J. F.; Gloaguen, F.; Schollhammer, P.; Talarmin, J. J. Electroanal. Chem. 2006, 595, 47−52. (50) Bourrez, M.; Steinmetz, R.; Gloaguen, F. Inorg. Chem. 2014, 53, 10667−10673. (51) Caplins, B. W.; Lomont, J. P.; Nguyen, S. C.; Harris, C. B. J. Phys. Chem. A 2014, 118, 11529−11540. (52) Mirmohades, M.; Pullen, S.; Stein, M.; Maji, S.; Ott, S.; Hammarstrom, L.; Lomoth, R. J. Am. Chem. Soc. 2014, 136, 17366− 17369. (53) Sicolo, S.; Bruschi, M.; Bertini, L.; Zampella, G.; Filippi, G.; Arrigoni, F.; De Gioia, L.; Greco, C. Int. J. Hydrogen Energy 2014, 39, 18565−18573. (54) Solis, B. H.; Hammes-Schiffer, S. Inorg. Chem. 2014, 53, 6427− 6443. (55) Raugei, S.; DuBois, D. L.; Rousseau, R.; Chen, S.; Ho, M.-H.; Bullock, R. M.; Dupuis, M. Acc. Chem. Res. 2015, 48, 248−255. (56) Bourrez, M.; Gloaguen, F. New J. Chem. 2015, 39, 8073−8079. (57) Lam, Y. C.; Nielsen, R. J.; Gray, H. B.; Goddard, W. A. ACS Catal. 2015, 5, 2521−2528. (58) Riplinger, C.; Carter, E. A. ACS Catal. 2015, 5, 900−908. (59) Rountree, E. S.; McCarthy, B. D.; Eisenhart, T. T.; Dempsey, J. L. Inorg. Chem. 2014, 53, 9983−10002. (60) Appel, A. M.; Helm, M. L. ACS Catal. 2014, 4, 630−633. (61) Artero, V.; Saveant, J.-M. Energy Environ. Sci. 2014, 7, 3808− 3814. (62) Felton, G. A. N.; Glass, R. S.; Lichtenberger, D. L.; Evans, D. H. Inorg. Chem. 2006, 45, 9181−9184. (63) Roberts, J. A. S.; Bullock, R. M. Inorg. Chem. 2013, 52, 3823− 3835. (64) Delahay, P.; Stiehl, G. L. J. Am. Chem. Soc. 1952, 74, 3500− 3505. (65) Nicholson, R. S. Anal. Chem. 1965, 37, 1351−1355. (66) Savéant, J.-M. Chem. Rev. 2008, 108, 2348−2378. (67) Costentin, C.; Drouet, S.; Robert, M.; Savéant, J.-M. J. Am. Chem. Soc. 2012, 134, 11235−11242. H

DOI: 10.1021/acs.inorgchem.5b02245 Inorg. Chem. XXXX, XXX, XXX−XXX

Forum Article

Inorganic Chemistry (100) Wang, H.-Y.; Wang, W.-G.; Si, G.; Wang, F.; Tung, C.-H.; Wu, L.-Z. Langmuir 2010, 26, 9766−9771.

I

DOI: 10.1021/acs.inorgchem.5b02245 Inorg. Chem. XXXX, XXX, XXX−XXX