Spectroscopic and Quantum Chemical Investigation of Benzene-1,2

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

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Spectroscopic and Quantum Chemical Investigation of Benzene-1,2dithiolate-Coordinated Diiron Complexes with Relevance to Dinitrogen Activation Sabrina I. Kalläne,† Anselm W. Hahn,† Thomas Weyhermüller,† Eckhard Bill,† Frank Neese,†,‡ Serena DeBeer,† and Maurice van Gastel*,†,‡ †

Max Planck Institute for Chemical Energy Conversion, Stiftstrasse 34-36, D-45470 Mülheim an der Ruhr, Germany Max-Planck-Institut für Kohlenforschung, Kaiser-Wilhelmplatz 1, D-45470 Mülheim an der Ruhr, Germany

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‡

S Supporting Information *

ABSTRACT: In this work, a benzene-1,2-dithiolate (bdt) pentamethylcyclopentadienyl di-iron complex [Cp*Fe(μ−η2:η4bdt)FeCp*] and its [Cp*Fe(bdt)(X)FeCp*] analogues (where X = N2H2, N2H3−, H−, NH2−, NHCH3−, or NO+) were investigated through spectroscopic and computational studies. These complexes are of relevance as model systems for dinitrogen activation in nitrogenase and share with its active site the presence of iron, sulfur ligands, and a very flexible electronic structure. On the basis of a combination of X-ray emission spectroscopy (XES), X-ray crystallography, Mössbauer, NMR, and EPR spectroscopy, the geometric and electronic structure of the series has been experimentally elucidated. All iron atoms were found to be in a local low-spin configuration. When no additional X ligand is bound, the bdt ligand is tilted and features a stabilizing π-interaction with one of the iron atoms. The number of lone-pair orbitals provided by the nitrogencontaining species is crucial to the overall electronic structure. When only one lone-pair is present and the iron atoms are bridged by one atom, a three-center bond occurs, and a direct Fe−Fe bond is absent. If the bridging atom provides two lonepairs, then an Fe−Fe bond is formed. A recurring theme for all ligands is σ-donation into the unoccupied eg manifolds of both iron atoms and back-donation from the t2g manifolds into the ligand π* orbitals. The latter results in a weakening of the double bond of the bound ligand, and in the case of NO+, it results in a weakening of all bonds that comprise triple bond. The electronrich thiolates further amplify this effect and can also serve as bases for proton binding. While the above observations have been made for the studied di-iron complexes, they may be of relevance for the active site in nitrogenase, where a similar N2 binding mode may occur allowing for the simultaneous weakening of the N2 σ bond and π bonds.

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INTRODUCTION Nitrogen fixation is one of the most energy demanding chemical energy conversion processes, owing to the high bond dissociation energy (BDE) of the dinitrogen triple bond of 226 kcal/mol.1 In nature, the binding of N2 and conversion to ammonia is accomplished by enzymes called nitrogenases according to the reaction N2 + 8H+ + 8e− → 2NH3 + H 2

heart of the enzyme contains a complex cluster consisting of an Fe7MoS9C (“FeMoco”) center,4,6 where binding of N2 and subsequent reduction to ammonia occurs via a series of intermediates dubbed E-states, in a mechanistic scheme proposed by Lowe and Thorneley.7−12 Moreover, an additional [4Fe4S] cluster and a formal [8Fe7S] cluster, called P cluster, are present,13 perhaps partaking by storing the reducing equivalents, without which the catalytic FeMoco cluster does not turn over. When the P cluster is present,

(1)

In this process, 16 ATP molecules are converted to ADP and one molecule of H2 is produced.2−5 As such, nitrogenases are able to perform an impressive eight-electron reduction. The © XXXX American Chemical Society

Received: January 21, 2019

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DOI: 10.1021/acs.inorgchem.9b00177 Inorg. Chem. XXXX, XXX, XXX−XXX

Article

Inorganic Chemistry

Figure 1. Schematic structures of complexes 1, 1+, 2, 2m, 2+, 3+, 3m+, 4+, 4m+, 5, and 6+ (Lut = 2,6-lutidine). In the text, the left iron atom is referred to as Fe1, and the right iron atom as Fe2. In 1 and 1+, bdt is bound asymmetrically. Both iron atoms σ interact with the bdt ligand, but only Fe1 is able to π interact with bdt. Complexes on the upper row are diferrous, FeII−FeII systems; complexes 1+, 2+, and 5 are paramagnetic, S = 1/2 species. Oxidation states as derived from experiments and calculations (vide infra) have been indicated. A formal Fe−Fe bond is present in 3+ and 4+ that feature a monatomic bridge (vide inf ra).

protons. Qu et al. carefully characterized the starting complex, [Cp*Fe(μ−η2: η4-bdt)FeCp*] 1, its singly oxidized complex 1+, complexes 2 and 2+ with bound diazene, asymmetric complex 3+ with bound N2H3−, and symmetric complex 4+ with a bound amido ligand (Figure 1) by crystallography and Mössbauer spectroscopy.18,19 Particularly relevant for the aspect of proton binding and triple bond activation are complexes 5 and herein newly reported complex 6+, with hydride and nitrosyl bridging ligands, respectively. We employ a combination of spectroscopy and theory, which in our view is indispensable for understanding the binding of the ligand and the weakening of the ligand bonds by electron-donation and back-donation mechanisms. Detailed studies of the electronic structure of these complexes can provide insight into how the storage of reduction equivalents is achieved, either in the form of a delocalized reduction equivalent that involves the Cp* and bdt ligands or in terms of two electrons that form an Fe−Fe bond. Of all complexes studied, new complex 6+ with a bridging nitrosyl ligand, NO+, which is isoelectronic to dinitrogen, is relevant as a model for investigating the electronic aspects of NN bond activation. Lastly, another interesting and so far not understood aspect that we address here is the observation that the bdt ligand in some of the complexes takes up a tilted orientation, which leads to a lowering the symmetry of the complex. We analyze these aspects, as well as the question of the low-spin versus high-spin multiplicity of the iron atoms and concomitantly whether or not the ligand may attain radical character, and put our findings into perspective with respect to activation of the NN triple bond.

FeMoco performs the monumental task of binding N2, activating the triple bond, harboring protons and reducing equivalents in the form of hydrides14,15 and being able to tune the free energies of the intermediates of all reduction steps, in order to ensure efficient catalysis. In industry, ammonia synthesis from N2 proceeds by the Haber−Bosch process in which a heterogeneous iron-based catalyst operates at typical pressures of 20 MPa and temperatures of 500 °C.16 One major difference between nature and the Haber−Bosch process is that nature is able to perform the reaction in a targeted way at ambient temperature and pressure. The Haber−Bosch process, however, presently happens on a scale of 450 million tons of fertilizer being produced per year, consuming 1% of the world’s energy supply. Hence, it would be advantageous to investigate the process of triple bond activation in N2 on a fundamental level, in order to perhaps 1 day rationally design catalysts for more efficient ammonia synthesis. While di-iron complexes that bind hydrazine has been reported already by Sellmann et al.,17 Li and Qu et al. synthesized and investigated a particularly relevant [Cp*Fe(μ−η2:η4-bdt)FeCp*] (bdt = benzyl-1,2-dithiolate) complex that reacts with an excess of hydrazine under formation of a diazene complex [Cp*Fe(μ−η2:η2-bdt)(μ−η1: η1-HNNH)FeCp*], and related complexes 1−5 (Figure 1).18−20 This reaction has similarities to the process of reduction of N2 to ammonia catalyzed by nitrogenase. Investigation of the electronic structures of the model complex with various bound ligands may yield information about aspects of NN bond activation and storage of reducing equivalents and B


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Inorganic Chemistry

Figure 2. (a) Mössbauer spectrum of 1 (black) and simulation (red) with two iron species (green and blue, both 50% abundance). (b) Iron Kβ mainline XES spectrum of 1. (c) Mössbauer spectrum of 1+(black) and simulation (red) with a majority species (green and blue, both 44%) and two minority species (magenta, 9% and yellow, 3%). (d) Iron Kβ mainline XES spectrum of 1+. (e) EPR spectrum (solid line) and simulation (dashed line) of 1+.

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C26H34BF4Fe2NOS2, M = 639.17, triclinic, space group P1̅, a = 10.0463(4) Å, b = 11.0514(8) Å, c = 12.7089(9) Å, α = 76.073(5)°, β = 86.095(4)°, γ = 84.461(5)°, V = 1361.67(15) Å3, ρcalcd. = 1.559 g cm−3, T = 100(2) K, Z = 2, μ(Mo Kα) = 1.267 mm−1, 25 366 reflections measured, 10 305 unique, (Rint = 0.0431). Final R1, ωR2 on all data: 0.0550, 0.1117. R1, ωR2 values for 8445 reflections with I0 > 2σ(I0): 0.0413, 0.1017; residual electron density +1.138, −1.565 e Å−3. The crystallographic data have been deposited with the Cambridge Crystallographic Data Centre as supplementary publication with a CCDC reference number 1874850. Spectroscopy. Mössbauer spectra were recorded on a conventional Mössbauer spectrometer with alternating constant acceleration of the γ-source. The minimum experimental line width was 0.24 mm/ s (full width at half-height). The sample temperature was maintained constant in an Oxford Instruments Variox cryostat. The γ-source (57Co/Rh, 1.8 GBq) was kept at room temperature. Isomer shifts are quoted relative to iron metal at 300 K. The spectra at zero field were collected for powder samples at 80 K. The spectra recorded were fitted using the program MFIT with Lorentzian doublets. X-band EPR spectra were recorded on Bruker Elexsys E500 (cw EPR) and E580 (pulsed ENDOR and ESEEM) spectrometers. For Kβ-emission measurements, all XES spectra were collected with an in-house X-ray emission spectrometer, which was designed based on a prototype of Anklamm et al.21,22 A liquid-gallium jet (Excillum) was utilized as an X-ray source and its Kα1 emission line at 9572 eV were used for the core hole ionization of Fe.23 The

EXPERIMENTAL SECTION 18

Synthesis. 1, 2, 2m, 2PF6, 3BPh4, 4BPh4, and 4mBPh4 as well as 1BF4 and 519 were prepared as described in literature. The complexes have been obtained typically in a purity of 85%. Although great effort has been undertaken especially with respect to oxidation of these very sensitive compounds, the presence of about 15% of unknown decay product, also present in the original studies18 could not be avoided. Synthesis of 6BF4. A solution of [Cp*Fe(μ−η2: η4-bdt)FeCp*] (1) (200 mg, 383 μmol) in dichloromethane (3 mL) was cooled to −60 °C and treated with a solution of NOBF4 (45 mg, 385 μmol) in dichloromethane (10 mL). The reaction mixture was allowed to warm slowly to room temperature and stirred for 1 h. The volatiles were removed under vacuum. The remaining brown solid was washed with THF (2 × 3 mL) and extracted with dichloromethane (5 mL). Addition of pentane to the solution results into the formation of a precipitate, which was separated and dried in vacuum. Yield: 168 mg (69%). Analytical Data for 6BF4. 1H NMR (500 MHz, CD2Cl2): δ = 7.40 (dd, 2H, J = 5 Hz, J = 5 Hz, CHbdt), 6.81 (dd, 2H, J = 5 Hz, J = 5 Hz, CHbdt), 1.51 (s, 30H, CH3) ppm. 19F NMR (470.5 MHz, CD2Cl2): δ = −153.9 ppm. FTIR (ATR): ν = 1579 (NO) cm−1. Brown crystals suitable for X-ray diffraction analysis were obtained by slow diffusion of pentane into a dichloromethane solution of 6+BF4‑. The diffraction data were collected on a Bruker KappaCCD diffractometer at 100 K for a fragment with the dimensions of 0.21 × 0.18 × 0.03 mm3. C


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Inorganic Chemistry

Figure 3. (a) Mössbauer spectrum of 2 (black) and simulation (red) with a majority species (green, 88%) and a minority species (blue, 12%). (b) Iron Kβ mainline XES spectrum of 2. (c) Mössbauer spectrum of 2+ (black) and simulation (red) with a majority species (green, 71%) and minority species (blue and magenta, 25% and 4%). (d) Iron Kβ mainline XES spectrum of 2+. (e) EPR spectrum (solid line) and simulation (dashed line) of 2+. The molecular z-axis is directed from the center of each Cp* to the respective iron atom; the local x-axes of both iron centers coincide and are chosen to be along the S−S direction. resulting Kβ Fe-emission were detected with a von Hamos type spectrometer.24 The emission energies were calibrated to a Fe2O3reference at the Kβ1,3 (7059.5 eV). The samples were prepared in Cusample holder as pressed pellets in an argon-filled glovebox directly connected to the instrument. Quantum Chemical Calculations. All electronic structure calculations reported in this work were performed with the ORCA program.25 Optimized model geometries are included in the Supporting Information. All calculations for singlet states are spinrestricted calculations. Those for states of higher spin-multiplicity and broken-symmetry calculations are spin-unrestricted calculations. Calculations are based on density functional theory (DFT) using the BP86 functional26 and scalar relativistic corrections included in form of the zero-order regular approach (ZORA).27,28 The matching ZORA-Def2-TZVP basis set29 was used for all calculations. The resolution of identity approximation30 was employed to speed up calculation of the two-electron integrals. Mössbauer parameters have been calculated by performing reference calculations on a set of molecules as used by Römelt et al.,31 which for the isomer shift δ and electron density at the iron atom ρ(Fe) gave rise to the following regression: δ [mm/s] = 3546.29−0.30617ρ(Fe). EPR g values have been calculated by linear response theory.32 Although a vast literature exists related to the ordering of spin states with varying functional, we rather prefer to use the experimental observation from NMR, XES and Mössbauer spectroscopy in order to establish whether a state is

diamagnetic. The magnetic nature of each complex as measured by NMR is specified in the Results section. The diagrams with orbital contours concern localized quasi-restricted orbitals. The SOMO in Figure 10 is formally called the singly occupied natural orbital.

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RESULTS AND DISCUSSION

In the Results section, we present Mössbauer, XES, and EPR spectra of compounds 1−6. Because we firmly believe that analysis of experimental data is best performed in combination with quantum chemical calculations, we refrain from interpretation until after presenting the quantum chemical data. Here, we limit ourselves to describing the spectra and the parameters extracted from the spectra by simulation and fitting procedures and solely indicate whether the iron atoms are equivalent. Spectroscopy. Complex 1 and 1+. Figure 2a displays the Mössbauer spectrum of 1, which features two quadrupole doublets in a 50:50 (±4%) ratio with an isomer shift of 0.43 and 0.63 mm/s, and quadrupole coupling constants of 2.13 and 1.96 mm/s, respectively (cf. Table 2). The sets of parameters differ significantly, which is representative of the inequivalence of the iron atoms, possibly induced by the tilt of D


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Inorganic Chemistry the bdt ligand. The Mössbauer spectrum is well-represented by a fit that features two quadrupole doublets in a 50:50 ratio. Figure 2b shows the iron Kβ mainline XES spectrum of 1. This spectrum corresponds to 3p to 1s transitions, and due to 3p−3d exchange in the final state, it serves as a sensitive probe of the iron local spin state.33,34 The lack of structure and absence of the Kβ1 feature in the spectrum suggests the iron sites are best described as ferrous, with S = 0 local spin configurations,35 similar to the Mössbauer data. The Mössbauer spectrum of oxidized compound 1+ contains multiple contributions which reveals that this compound was not obtained with 100% purity. The isomer shifts of 0.43 and 0.57 mm/s and quadrupole splittings of 0.57 and 1.19 mm/s do confirm the presence of two inequivalent iron atoms. The fit indicates that the yield of 1+ upon oxidation of 1 amounts to 88% abundancy in with both irons contributing equally, i.e., in a 44:44% ratio (Figure 2c). In the XES measurements of 1+, the characteristic structureless Kβ′ mainline spectrum prevails, suggesting that the irons still exist in low-spin electronic configurations (Figure 2d). The Kβ1,3 maximum at 7058.7 eV remained constant in emission energy during the oxidation, in agreement with previous observations on related systems, which have shown that the contributions of a change in oxidation state may be countered by changes in covalency.36,37 The majority species in the EPR spectrum of 1+ (Figure 2e) is characterized and well-simulated by g values of 3.14, 1.96, and 1.93. A minority species has been included in the fit as well, with g values of 4.47, 2.32, and 1.88. The g values of the majority species support the presence of a low-spin, S = 1/2, ferric iron. Complex 2 and 2+. In contrast to 1, the Mössbauer spectrum of 2 (Figure 3a) displays only one quadrupole doublet and confirms the equivalence of both iron atoms. This is in good agreement with the C2 symmetry of 2. Isomer shifts and quadrupole splittings of the majority species (88% abundance, cf. Figure 3) of 0.43 and 1.45 mm/s, respectively, fully supporting the low-spin FeII nature. Similar to 1, the Kβ XES spectrum of 2 retains a featureless mainline and supports a local low-spin electronic configuration at both irons (Figure 3b). The Mössbauer spectrum of 2+ (Figure 3c) displays one doublet and features a majority species with an isomer shift of 0.40 mm/s and a quadrupole splitting of 0.73 mm/s, in agreement with the presence of two equivalent iron atoms and C2 symmetry of the complex. The XES spectrum (Figure 3d) retains its shape, indicating the iron atoms invariably remain low spin. Complex 2+ is paramagnetic. The EPR spectrum (Figure 3e) displays one signal with g values equal to g|| = 2.85 and g⊥ = 1.97. The g shifts are markedly smaller than those of 1+ (cf. Figure 2e). Since the g shift is proportional to the spin population at iron, it provides a first hint that the unpaired electron in 2+ may be more delocalized than in 1+. Complex 3+. The Mössbauer spectrum of 3+ (Figure 4) gives rise to main species (85% abundance) with two distinct quadrupole doublets and isomer shifts of the two main species of 0.43 and 0.51 mm/s and quadrupole splittings of 0.94 and 1.22 mm/s, respectively. As before, the small isomer shifts indicate local low-spin configurations of the Fe3+ atoms. The fit of the Mössbauer spectrum additionally reveals the presence of an unknown minority species with 15% abundance.

Figure 4. Mössbauer spectrum of 3+ (black) and fit (red) consisting of two main species (blue and green) of 85% abundance and minority species (yellow, magenta) of 15% abundance.

Complex 4+. The Mössbauer spectrum of 4+ is shown in Figure 5. The spectrum features one quadrupole doublet,

Figure 5. Mössbauer spectrum of 4+ (black) and fit (red) consisting of a majority species (green, 94%) and a minority species (blue, 6%).

indicative of two equivalent iron atoms. The isomer shift of 4+ in this respect is similar to that of 2 and amounts to 0.42 mm/ s (Figure 5). However, the quadrupole splitting increases drastically to 2.78 mm/s (2.74 mm/s for 4m+, data not shown), as compared to 1.45 mm/s for 2. Complex 5. The Mössbauer spectrum of 5 (cf. Figure 6a) displays a majority species (85% abundance) with one quadrupole doublet. The isomer shift and quadrupole splitting parameters are 0.44 and 1.28 mm/s (cf. Table 2), again indicating the equivalence of both low-spin iron atoms. The EPR spectrum is displayed in Figure 6b and the majority species is characterized by g values of 2.51, 2.07 and 2.00, confirming that 5 is a paramagnetic, S = 1/2 system. Since the hydride in 5 is directly coordinated to the iron centers, additional experiments have been performed with the aim of resolving the hyperfine coupling constants of the hydride. The rationale for performing these experiments is that the hyperfine tensor of the hydride provides very sensitive local information about the spin-density distribution. It provides insight into whether a metal−metal bond exists, and whether the hydride 1s orbital participates in the SOMO or in a lower-lying, doubly occupied orbital. Two complementary methods have been used: the electron spin echo envelope modulation (ESEEM), which is most sensitive for E


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Inorganic Chemistry

Figure 6. (a) Mössbauer spectrum of 5 (black) and fit of one main species (green) with a weight of 85% and two minority species (blue, magenta) of 8% and 7% abundance; (b) continuous-wave (cw) EPR spectrum (10 K, 9.63 GHz, microwave power 2 mW, modulation amplitude 0.5 mT) of 5 in dichloromethane. Simulations are included (dashed line). The g-values amount to 2.50, 2.07, and 2.00 for 5. (c) X-band Davies ENDOR and 3-pulse ESEEM spectra (10 K, 9.715 GHz) of 5. Length of 90 pulse: 16 ns (ESEEM) or 80 ns (ENDOR); length of RF pulse: 8 μs; time between first and second pulse (ESEEM): 200 ns. Simulations are given in dashed lines. The ENDOR experiment is more sensitive at large radio frequencies (>15 MHz), whereas ESEEM is more sensitive in the lower frequency region, but not at canonical orientations. These intensity effects have been neglected in the present simulations, since the emphasis was on reproducing the positions of the bands. F


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Inorganic Chemistry

0.03°), which does point to the presence of a bound NO+ and ferrous iron atoms rather that NO− and ferric iron atoms. The FTIR spectrum displays an N−O stretching frequency of 1579 cm−1, much lower than that of free NO+ (2250 cm−1) and also lower than that of free NO (1876 cm−1),39 but larger than that of NO− (1284 cm−1). The frequency is in a range typical for a bridging NO ligand with a formal positive charge and comparable with that of [(LFe)2NO]BF4 (1553 cm−1, L = N,N′-dimethyl-N,N′-bis(β-mercaptoethyl)ethylenediamine).40 Quantum Chemistry. Complex 1. The crystal structure, as well as the NMR and IR data of the diamagnetic [Cp*Fe(μ−η2:η4-bdt)FeCp*] complex 1 have been previously reported by Li et al.18 Selected bond distances and angles from the crystal structure are given in Table 1. A structural feature of particular interest is the asymmetric coordination of the bdt ligand. Specifically, the Fe1−S distances are elongated by 0.14 Å as compared to the Fe2−S distances and the Fe1−S−Cbdt angle amounts to only 61°, thus positioning the C1 and C2 atoms of the bdt ligand at only 2.10 Å distance from Fe1 and 3.16 Å from Fe2 (see also Figure 1). Previously, Li et al. attributed this asymmetry to the presence of a π-radical monoanionic bdt ligand by analogy to work by Ray et al.41 This would result in a formal oxidation state assignment of Fe(I)Fe(II) in complex 1. DFT relaxed surface scans of the Fe1−S−C1 angle, i.e., a geometry optimization in which all coordinates except the Fe1−S−C1 angle are optimized, have been performed (Figure 8a) in order to elucidate the origin of the tilt of the bdt ligand. Moreover, all possible spin couplings have been considered for the putative case of both iron atoms being high-spin, S = 2, which would in principle afford one nonet (S = 4), one septet (S = 3), one quintet (S = 2), one triplet (S = 1), and one singlet (S = 0) state. The relaxed surface scan indicates that the singlet state is lowest in energy, with the nonet state no less than 60 kcal/mol above the ground state. Moreover, attempts to obtain broken symmetry solutions that would be indicative of a diradical species invariably converged to a closed-shell singlet state with local low-spin, S = 0, ferrous ions. Hence, in the calculation there is no bdt radical anion character. The singlet surface displays two identical minima, which derives from the symmetry of the complex in that the bdt can tilt to either Fe1 or Fe2. The tilted geometry of the bdt ligand turns out to be favored by about 15 kcal/mol over the geometry with a symmetrically oriented bdt. The symmetric orientation of the bdt ligand is only compatible

nuclear transitions close to 0 MHz and up to 6 MHz at Xband microwave frequencies;38 and electron nuclear double resonance (ENDOR), which typically becomes the more sensitive method for nuclear frequencies above 6 MHz. The recorded set of orientation-selected ESEEM and ENDOR spectra (Figure 6c) display well-resolved signals with bands up to 37 MHz. Simulation of the positions of the bands in the ESEEM and ENDOR signals (Figure 6c, dashed line) has been achieved with a program written in-house and by reading the effective values from those spectra that correspond to the canonical orientations as starting parameters. The principal values of the hyperfine tensor amount to 14, 24, and 44 MHz (±1 MHz). Note that the sign of these couplings cannot be determined from experiment. Complex 6+. Lastly, NOBF4 has been added to 1, affording new, symmetric, pure, and diamagnetic complex 6+. It has Mössbauer isomer shifts and quadrupole parameters of 0.40 and 2.09 mm/s, respectively (Figure 7). The complex has been

Figure 7. Mössbauer spectrum (black) of 6+ and fit (red).

crystallized, from which it becomes clear that the complex features a C2 symmetry axis. The full crystal structure is provided in the Supporting Information. We mention here that the crystal structure features an Fe−Fe distance of 2.385 Å which is shorter than that of 5 (Table 2). Moreover, the two iron atoms and the nitrogen and oxygen atoms of the NO ligand are all located in the same plane (dihedral angle of

Table 1. Selected Bond Lengths [Å] and Angles [°] for 1, 2m, 2PF6, 3BPh4, 4BPh4 and 4mBPh4,18 1BF4, 5,19 and 6BF4

Fe1−Fe2 Fe1−S Fe2−S Fe1−Cbdt Cp*−Fe1 Cp*−Fe2 Fe1−S−Cbdt Fe1−N Fe2−N Fe1−H Fe2−H N−N N−O

1

1BF4

2m

2PF6

4BPh4

4mBPh4

5

6BF4

2.76 2.29 2.15 2.10 1.67 1.66 61

2.68 2.27 2.17 2.14 1.71 1.73 63

3.22 2.34 2.34

3.18 2.34 2.34

2.70 2.23 2.23

3BPh4

2.44 2.22/2.33 2.22/2.33

2.44 2.21/2.34 2.21/2.34

2.42 2.27 2.26

2.3851(3) 2.2334(5)/2.3449(5) 2.2385(5)/2.3448(5)

1.72 1.72

1.73 1.74

1.71 1.75

1.72 1.72

1.73 1.73

1.70 1.70

1.72 1.72

1.87 1.87

1.83 1.83

1.89/1.94 1.91

1.93 1.93

1.90 1.90

1.826(1) 1.813(2) 1.58 1.62

1.29

1.31

1.40 1.207(2) G


Article

Inorganic Chemistry

Figure 8. (a) Relaxed surface scan of the Fe1−S−C1 bond angle in 1, performed for spin multiplicities S = 0 up to S = 4. (b) 3d orbital manifolds of both ferrous irons in 1; the HOMO is located at Fe2, a π bond is present between the dxy orbital of Fe1 and the C1−C2 π* orbital. The local zaxis is directed from the center of each Cp* to the respective iron atom, the local x-axes coincide and are chosen to be along the S−S direction.

significantly much longer Cp*−-Fe bond length of 1.97 Å.42 A complete list of Cartesian coordinates for the geometry optimized structure is provided in the Supporting Information. Analysis of the orbital structure quickly reveals the origin of the tilt (Figure 8b). Rather than the presence of a bdt radical, most notably, a stabilizing π-interaction is present between the dxy orbital of Fe1 and the bdt C1 and C2 carbon atoms (Figure 8b, lower-left orbital). While there is in principle no preference for a particular iron due to the symmetry of the molecule, the magnitude of the stabilization induced by the π interaction (15

with a triplet ground state (Figure 8a), which is not compatible with the diamagnetic NMR spectrum. The calculated Fe−Fe distance for the geometry-optimized singlet state is 2.83 Å, in reasonable agreement with the experimental value of 2.76 Å. The calculated Fe1−S−Cbdt angle is 61 degrees, in excellent agreement with experiment (cf. Table 1). Moreover, the distance from Fe1 and Fe2 to the center of the respective Cp* ligands of 1.67 and 1.66 Å (cf. Table 1), respectively, suggests a local low-spin character of the iron atoms, because high-spin Fe2+ complexes display a H


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Inorganic Chemistry

effectively oxidized by 25% of an oxidation equivalent. Together, they carry 50% of the oxidation. The bdt-phenyl moiety carries about 20%. Thus, there is a significant electronic relaxation taking part by the oxidation of Fe2. The calculated Mössbauer parameters are δ = 0.68 and ΔEQ = 1.45 mm/s for Fe1 and δ = 0.51 and ΔEQ = 1.18 mm/s for Fe2. The quadrupole splitting for Fe2 is thus not reproduced well for 1+, and the quadrupole splitting for Fe1 is calculated too small by 0.3 mm/s (cf. Table 1). Still, although quantitative agreement is not reached, both the calculated and the experimental quadrupole splittings decrease for both iron atoms when going from 1 to 1+, and in particular for Fe2, which is the center of oxidation. The calculated Mössbauer shifts are at least able to reproduce the trends in the isomer shift and quadrupole coupling. Taking all experimental and computational data together, 1+ is best described as a low-spin, localized valence FeII−FeIII complex with bdt tilted toward the ferrous iron, because it can more efficiently back-bond than the ferric site. Complex 2. After binding of a bridging N2H2 ligand complex 2 is formed. The crystal structure of Li et al.18 indicates that 2 has C2 symmetry with the bdt ligand in a symmetric conformation (see Figure 1). As before, the short Fe−Cp* distance (e.g., 1.72 Å for 2m, Table 2) suggests the presence of two low-spin ferrous iron atoms. The Fe−Fe distance amounts to 3.22 Å, making the presence of an Fe−Fe bond unlikely. An NMR spectrum has also been reported confirming the diamagnetic nature of 2.18 The calculated geometric parameters all agree with the crystal structure. DFT calculations have been performed with the aim to investigate the origin of the symmetric orientation of bdt. Full geometry optimization for the singlet state indeed leads to a symmetric orientation of the bdt ligand. The frontier orbitals of both ferrous irons are shown in Figure 9, left. The orbitals are pairwise degenerate and, as expected, are made up of the 3dz2, 3dx2−y2, and 3dxy pairs, respectively. The 3dxz and 3dyz orbitals are unoccupied. Interestingly, a new stabilizing interaction in 2 is present between the dxy orbitals of both Fe ions and the N2H2 π* orbital. This new π back-donation occurs from both iron atoms, stabilizes the complex and also weakens the NN bond (calculated NN bond length 1.29 Å, see Table 2, as opposed to a standard NN bond length of 1.25).1 Moreover, it is clear from inspection of the lower lying orbitals that the N2H2 unit is a neutral diazene ligand and not a dianionic hydrazido ligand, since two doubly occupied bonding orbitals, one σ and one π orbital, have been found between the nitrogen atoms. The NN π orbital and the N lone-pair orbitals are included in Figure 9, right. Calculated Mössbauer parameters of 2 amount to 0.51 mm/ s for the isomer shift and 1.21 mm/s for the quadrupole splitting. The calculated Fe−Fe distance is 3.26 Å. The calculated parameters agree with their spectroscopically determined counterparts. Complex 2+. Oxidation affords complex 2+ for which a crystal structure exists as well.18 Selected geometric parameters are included in Table 2. Complex 2+ still features a symmetric bdt ligand and apart from minor changes in bond length, largely retains the same structure as that of 2. As such, the symmetric structure of the complex implies both iron atoms are equivalent, hinting to a mixed-valence, Fe2.5Fe2.5 complex. Our calculations indicate that oxidation from 2 to 2+ allows for increased σ donation of the N lone pairs to iron, which is

kcal/mol) keeps bdt tilted to either side, and disfavors a symmetrically oriented bdt, where neither of the iron dxy orbitals are stabilized. The stabilization of Fe1 additionally results in the dx2−y2 orbital of the other iron (Fe2) being the HOMO of 1 (Figure 8b, upper right orbital). In fact, close inspection of the crystallographic data from the work of Li et al.18 provides an a posteriori experimental indication of the π-interaction between the dxy orbital and the C1−C2 π* orbital. Specifically, the experimentally found C1− C2 bond length in the bdt elongates by about 0.015(8) Å in the tilted bdt structure (structure 1a of Li et al),18 as compared to a structure with a symmetrically oriented bdt (structure 2b of Li et al).18 A π-radical, if present, would rather feature a shortened C−S bond length. The calculated isomer shifts for 1 amount to 0.43 and 0.67 mm/s; the quadrupole splittings equal 1.97 and 1.88 mm/s and are in good agreement with experiment (Table 2), thus providing additional validation for the correctness of the electronic structure of 1 as presented in Figure 8b. Table 2. Mössbauer Parameters (Isomer Shift δ and Quadrupole Splitting |ΔEQ|) of All Investigated Complexes as Well as Calculated Valuesa δ [mm/s]

|ΔEQ| [mm/s]

complex

exp

calcd

exp

calcd

1 1BF4 2 2m 2PF6 3BPh4 4BPh4 4mBPh4 5 6BF4

0.43/0.63 0.43/0.57 0.43 0.47 0.40 0.44/0.51 0.42 0.42 0.43 0.40

0.54/0.73* 0.51/0.68* 0.51 0.50/0.57 0.44 0.48/0.59 0.51 0.52 0.52 0.51

2.13/1.96 0.57/1.19 1.45 1.45 0.73 0.98/1.23 2.78 2.74 1.27 2.09

1.97/1.88* 1.18/1.45* 1.21 1.22/1.41 0.76/0.72 0.82/1.44 2.60 2.71/2.72 1.25 2.17

a

Calculated values with asterisk apply to Fe1.

Complex 1+. Upon oxidation of 1, the main product 1+ is a paramagnetic, S = 1/2, species. The structural data of 1[BF4]19 reveal a largely unchanged metal cluster with a tilted bdt ligand. The most notable difference is a shortened Fe1−Fe2 distance of 2.68 Å in 1+ as compared to 2.76 Å for 1 (Table 1). The Cp*−Fe distances measured from Fe to the center of the Cp* ring amount to 1.71 and 1.73 Å, respectively, and are still representative of low-spin configurations of both iron atoms. The bdt C−C distance of 1+ amounts to 1.406(5) Å,19 significantly elongated as compared to existing data for 2+ (1.394(5) Å).18 DFT calculations confirm that the singly occupied molecular orbital (SOMO) of 1+ involves dx2−y2 character at Fe2. The SOMO of 1+ is essentially equal to the HOMO of 1 (Figure 8b). All doubly occupied orbitals of 1+ also do not change much as compared to those of 1, which is in agreement with the small changes observed in the crystal structure and XES spectra. In as far as the actual change in charge distribution is concerned, the formal oxidation of FeII to FeIII actually leads to a change in the Löwdin charge of Fe2 of only 0.15 oxidation equivalents when going from 1 and 1+. The electronic structure reveals that another 10% is carried by both sulfur atoms. Fe1 carries 3% of the oxidation, each Cp* gets I


Article

Inorganic Chemistry

Figure 9. (Left) 3d orbital manifolds of both ferrous irons in 2; the orbitals are pairwise degenerate, a π bond is present between the dxy orbitals of Fe1 and Fe2 and the NN π* orbital; also indicated (right) are the doubly occupied NN π orbital and the respective N lone pair orbitals.

significantly smaller than that observed in experiment.32 The calculated Mössbauer parameters for 2+ (Table 2) agree well with experiment. Complex 3+. Protonation and subsequent intramolecular electron transfer by addition of a proton source to 2 affords 3+ and leads to a change of the hapticity of the bdt ligand, as well as a rearrangement of the nitrogen containing ligand. According to the crystal structure,18 the bdt ligand features one bridging and one nonbridging thiolate. Moreover, NMR studies with 15N labeled substrates show that protonation of the HNNH ligand triggers electron transfer from the iron atoms, leading to the presence of two ferric ions and a N2H3− ligand.18 As such, the second bridging position is taken by the NH− fragment. As expected, the N−N bond elongates significantly to 1.45 Å upon protonation. The complex thus features a classical N−N single bond, where the experimentally found distance is slightly shorter than the distance of 1.50 Å calculated for an isolated N2H3−. The slight shortening likely happens owing to the presence of the positively charged iron atom. The calculated geometry (see the Supporting Information) again agrees with the crystal structure.18 For completeness, the state with maximum possible spin multiplicity, S = 4, is calculated to be 53 kcal/mol higher in energy than the singlet ground state. The localized 3d orbitals for the N2H3− complex 3+ of both Fe 3d manifolds as well as the relevant doubly occupied ligand orbitals are shown in Figure 11. Since N2H3− features an N−N single bond, no π* acceptor orbital for back-donation is available. This is indeed reflected in the localized orbital scheme. Complex 3+ features a true Fe−Fe bond formed by the σ combination of the dx2−y2 orbitals. These orbitals overlap, owing to the short Fe−Fe distance and the much larger tilt of the local molecular axes of each iron, also exemplified by the fact that the planes of both Cp* moieties are much less parallel

concomitant with increased Fe 3d character into N the lonepair orbitals. The increased σ donation strengthens the Fe−N bonds and weakens the NN bond and thus shortens the Fe−N bond distances from 1.87 to 1.83 Å (cf. Table 2). Additionally, since oxidation removes an electron from the dxy set, the amount of π-backbonding is reduced. The latter causes a shortening of the NN bond, so the net effect of increased σ donation and reduced π back-donation leads to similar N N bond lengths in 2 and 2+. Other than the minor but important change to the σ donation of the nitrogen lone-pair orbitals in the doubly occupied orbitals, leading to small geometric changes mentioned in the previous paragraph as compared to 2, the SOMO of 2+ is found to be evenly delocalized over both iron atoms, fully compatible with a mixed-valence Fe2.5Fe2.5 complex (Figure 10). The calculated g values of g|| = 2.25 and g⊥ = 2.03 are certainly compatible with experiment, given the fact that the largest g value is known to be calculated systematically and

Figure 10. Singly occupied molecular orbital of 2+. J


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Inorganic Chemistry

Fe−Fe bond direction as in complex 2. However, the Fe−Fe distance shortens to 2.44 Å and the Fe−N distances amount to 1.93 Å for both 4+ and 4m+. The complexes are diamagnetic, as evidenced by the presence of an NMR spectrum.18 Calculated bond distances (Fe−Fe: 2.42 Å; Fe− N: 1.91 Å) agree well with experiment. Investigation of the electronic structure of 4+ gives rise to a bonding pattern very similar to that of 3+; the relevant frontier orbitals and ligand orbitals are shown in Figure 12. An Fe−Fe

Figure 11. 3d orbital manifolds of both low-spin ferric irons and relevant ligand orbitals in 3+; a metal−metal bond is present, formed by the σ combination of the respective dx2−y2 orbitals. Also included are the three Fe−N bonding orbitals.

in 3+ than in those 2+. In total, 10 electrons are present in the 3d manifolds of both irons, and two of these are associated with the Fe−Fe bond and shared between both iron atoms. Although no unpaired electrons are present, the complex is best described as a closed-shell, diamagnetic complex harboring two low-spin ferric iron atoms that form a 2c2e Fe−Fe bond. None of the orbitals of both Fe 3d manifolds strongly interacts with N2H3−. The nitrogen atoms act as σ donors (Figure 11, bottom). The calculations indicate the presence of two lone-pair orbitals for the NH− fragment and one for the NH2 fragment with a small amount of Fe character mixed in, thus formally confirming the presence of an N2H3− ligand, i.e., 14 valence electrons, divided over 3 N−H, 1 N−N and 3 N− Fe bonding orbitals, of which the latter are included in Figure 11. Calculated Mössbauer parameters (isomer shifts 0.51 and 0.68 mm/s and quadrupole splittings 1.18 and 1.45 mm/s, cf. Table 2) agree well with experiment. Complex 4+. Subsequently, complex 4+ (or 4m+) with a bridging amido (or methylamido) ligand has been investigated. In both cases, the crystal structures18 shows that the bdt ligand attains a symmetric orientation perpendicular to the

Figure 12. 3d orbital manifolds of both low-spin ferric irons and relevant ligand orbitals for 4+; a metal−metal bond is present, formed by the σ combination of mainly the respective dz2 orbitals at both iron atoms.

bond is present in this complex as well. Two lone pairs of the amido ligand are present in the lower-lying bonding manifold that σ-donate into the iron 3d orbital set, thus overall classifying the ligand as NH2−, i.e., 8 valence electrons divided over 2 N−H and 2 N−Fe orbitals, the latter are included in Figure 12, and the iron atoms as a diamagnetic diferric pair that features an Fe−Fe bond, similar to 3+. The calculated Mössbauer parameters (isomer shift: 0.52 mm/s; quadrupole splitting: 2.71 and 2.72 mm/s) are included in Table 2. They agree well with their experiment counterparts. Complex 5. In the crystal structure of paramagnetic hydride complex 5, the bdt ligand attains a symmetric orientation.18 The Fe−Fe distance reduces to 2.42 Å. The Fe−H distances K


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Inorganic Chemistry are 1.58 and 1.62 Å.18 The number of signals in the paramagnetic NMR spectrum provided further evidence that the two Fe sites are equivalent.18 As before, the optimized geometry of 5 (Cartesian coordinates provided in the Supporting Information) agrees with the crystal structure. With the bdt ligand being symmetrically oriented, the unpaired electron is expected to be associated with the dxy orbital set, essentially identical to the case of 2+. Indeed, analysis of the orbital structure gives rise to two nearly degenerate linear combinations of the dxy orbitals as SOMO−1 and SOMO (Figure 13) and an evenly distributed spin density in which the contours of the dxy orbitals can be clearly recognized (Figure 13, dashed box). The interaction of the hydrido ligand involves a lower lying, bonding orbital, which

mainly involves the 1s orbital with a small admixture of the respective dyz orbitals (Figure 13, bottom). This doubly occupied orbital confirms that the hydride bridge is formally indeed H−. The orbital could in fact also be considered a three-center two-electron (3c2e) bonding orbital, although the hydride part largely dominates. A direct Fe−Fe bond was not found. However, the aforementioned doubly occupied delocalized (dxy)π bonding orbital and the SOMO could be considered as half a very weak π bond. With 11 electrons in the two 3d orbital manifolds and the unpaired electron being equally delocalized over both iron atoms (Figure 13), the complex features an Fe2.5Fe2.5 core and clearly lacks an Fe−Fe σ bond. The calculated g values are 2.22, 2.06, and 2.01, again in agreement with experiment. As for the hyperfine tensor of the hydride, calculated principal values of −6, −18, and −39 MHz are found (the principal axes of the g tensor and of hydride hyperfine tensor are nearly colinear), in excellent agreement with the absolute values deduced from the ENDOR and ESEEM spectra. In particular, the hydride hyperfine coupling constants, along with the calculated Mössbauer parameters that also agree with experiment, gives great confidence that the structural models, the local low-spin electronic configurations of the iron atoms and the formalisms used to describe the electronic structure are adequate. Complex 6+. The calculated Fe1−Fe2 distance of 6+ amounts to 2.389 Å. The calculated Fe1−S and Fe2−S distances equal 2.24 and 2.34 Å, whereby the sulfur atom opposite to the NO fragment (O−N−S angle: 162°) has the shorter distance to both iron atoms. The NO ligand (calculated N−O distance 1.207 Å; experimental distance: 1.207 Å) forms a plane with both iron atoms (calculated dihedral angle 0.8°), and the bdt ligand takes a symmetric orientation. All calculated geometric parameters agree well with the crystal structure of 6+. Complex 6+ is diamagnetic, and as before, a closed-shell and a broken symmetry calculation have been performed, of which both converged to a closed-shell wave function. The orbital structure of nitrosyl bound complex 6+ is presented in Figure 14. Of particular interest here is the question of whether NO carries a formal positive or negative charge, whether it may even be charge-neutral, the latter being unlikely since no broken symmetry solution was found. Inspection of the orbitals in Figure 14 provides direct support for the presence of NO+. First, in the iron 3d manifold, 6 doubly occupied orbitals that mainly have Fe 3d character have been found. As opposed to the previous complexes, 4 out of the 6 orbitals, i.e., two for each respective iron, feature significant back-donation into nitrogen pz and px orbitals (the molecular y direction is approximately along the NO bond). At lower cutoff values, the contours of the orbitals show that these orbitals are π-antibonding to oxygen, i.e., they are ligand π* orbitals. The presence of two π* orbitals in NO is only compatible with a triple bond, which in turn is only compatible with the presence of an NO+ ligand. Second, the set of ligand orbitals features 5 valence orbitals centered on NO; they comprise one σ and two π orbitals, one lone pair at O, and a lone pair at N, which σ interacts with both iron atoms. Thus, with 5 doubly occupied valence orbitals, i.e., 10 valence electrons, the calculation thus very clearly supports the presence of a formal NO+ ligand. Upon examination of the charge distribution, the Fe−N π back-donating orbitals are found to feature electron donation

Figure 13. 3d orbital manifolds of both irons ions and the hydride 1s orbital for 5; a metal−metal interaction is present as a doubly occupied dxy-based π orbital. The SOMO is given by the dxy-based π* orbital, thus classifying the formal oxidation states of both iron atoms at 2.5. L


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Inorganic Chemistry

symmetric position. In the case of complex 2 with coordinated diazine, N2H2, σ-donation of the nitrogen lone pairs to iron weakens the NN σ bond, and π-back-donation into the π* orbital of N2H2 also weakens the NN π bond. Upon oxidation to 2+, σ-donation is increased, thus further weakening the NN σ-bond, but π-back-donation decreases, strengthening the N−N π bond. For N2H3−, the geometry of complex 3+ changes to an asymmetric coordination, whereby the NH− fragment takes a bridging position and the NH2 fragment coordinates to only one iron atom. This demonstrates the flexibility of the [Cp*Fe(bdt)FeCp*] complex, whereby the geometric and electronic structures adjust in order to accommodate σdonation of all three nitrogen based lone-pair orbitals, i.e., two for the bridging NH− fragment and one for the NH2 fragment, altogether leading to a weakening of the N−N σ-bond. Analysis of simpler complexes 4+ and 5 with coordinated NH2− and H− ligands, respectively, indicate that the [Cp*Fe(bdt)FeCp*] framework even allows sufficient flexibility for storing reducing equivalents, be it either in the confines of a formal Fe−Fe bond in the former case or directly at the iron atoms in the latter case. Additionally, the mere fact that the complex allows the coordination of a hydride shows that it can also harbor protons. The propensity to store reducing equivalents in the Fe2 motif as well as the Lewis basic functionality for harboring protons, or equivalently formulated, the Lewis acidic functionality for hydride binding, in close proximity to the reducing equivalents is an extremely important aspect for catalytic reduction of N2 to NH3. In this respect the bdt ligand helps since both sulfides are electron-rich σ and π donors to iron that increase the propensity of iron for backbonding into ligand π* orbitals. The most spectacular observation relevant for N2 activation comes from analysis of the new complex, 6+, which features a bound NO+. NO+ is isoelectronic with N2. Upon coordination of NO+ to [Cp*Fe(bdt)FeCp*], the lone-pair of N is able to σ donate into empty 3d orbitals at both iron atoms, thus weakening the NO σ-bond. Moreover, the ferrous iron atoms of 6+ are able to π-back-donate into both NO π* orbitals, thus also weakening both π bonds. Hence, by being able to weaken all three elements of the triple bond of NO+, the [Cp*Fe(bdt)FeCp*] moiety seems to capture an important electronic aspect of activating a triple bond. This observation may well be relevant for the initial activation of the triple bond of the isoelectronic N2 by nitrogenases. For the mere binding of N2, however, it suffices that one of the lone pairs of N2 is able σ-donate into an unoccupied d orbital, most prominently the dz2 orbital for terminal binding of N2 to a single metal. We last address the question of spin-multiplicity of the iron atoms in relation to NN activation. Even though FeMoco features local high-spin configurations of the iron atoms, one would think that in particular low-spin ferrous iron should be well-suited for N2 activation for the simple reason that the empty eg set would allow optimal σ-donation from a ligandbased lone-pair and that the doubly occupied t2g set is ideally suited for π-backbonding into one or even two empty π* orbitals of the ligand. Such situation indeed seems to occur in 6+. If iron is in the ferrous, high-spin (S = 2) configuration, then all d orbitals carry at least one electron, so both σdonation from a ligand orbital is reduced and π back-donation will be less pronounced as well. In order to validate this hypothesis, the calculation of complex 6+ in the nonet (S = 4)

Figure 14. 3d orbital manifolds of both irons ions as well as the 5 valence orbitals of the NO ligand for 6+, pictorially representing the resonance structure of NO+ with one σ and two π orbitals. Also indicated for completeness is the resonance structure of NO+.

from iron to NO that significantly lessens the formally positive charge at NO to an actual much smaller than 1+. No Fe−Fe bond has been found, despite the short Fe−Fe distance of 2.385 Å. Since complex 6+ has no unpaired electrons, an assignment of formal oxidation states in terms of two ferrous iron atoms and NO+ thus provides a good description of the charge distribution in this complex. The calculated Mössbauer parameters agree with the experiment data (cf. Table 2). The calculated frequency for the NO stretching mode of 1569 cm−1 differs only by 10 cm−1 from the experimental value. As a final note, a B3LYP calculation gave rise to an NO stretching mode at 1637 cm−1. Although this number differs by 58 cm−1 from the BP86 calculation, the B3LYP functional qualitatively gives rise to an identical electronic structure as the BP86 functional. Summary and Relation to NN Bond Activation. This work represents an investigation of the electronic structure of the [Cp*Fe(bdt)FeCp*] complex with a set of ligands that invariably features iron atoms in their low-spin oxidation states. The complexes feature rich flexibility in the electronic structure with respect to ligand binding. The diiron complex in the case of 6+ allows for simultaneous π-backdonation into both ligand π* orbitals of NO+. Even the interaction with the bdt ligand is unusual in that it has the propensity of stabilizing one of the doubly occupied iron 3d orbitals (the 3dxy orbital in Figure 2) by tilting toward either iron atom and allowing π-back-donation into an empty bdt π* orbital. Upon interacting with nitrogen-based small molecules, a stabilization of the 3d manifold is accomplished through a larger stabilizing π-back-donation with the bound small molecule rather than with bdt, which accordingly assumes a M


Inorganic Chemistry

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state, i.e., with both ferrous ions being high-spin and ferromagnetically coupled, has been examined. This state is calculated to be 76 kcal/mol less stable than the singlet state. While σ-donation of the ligand lone-pair orbitals in this calculation still occurs, it seems to involve the spin-down electrons of the ligand, thus giving rise to the presence of 30% spin density at the NO + fragment. Additionally NO + coordination occurs with a much more a bent geometry, whereas in the singlet state, the iron atoms and the nitrogen and oxygen atoms are approximately in one plane. Although the energy of such high-spin state is significantly high above the singlet ground state in the molecule under investigation, it does give rise to the notion that ferrous iron in the high-spin state could in principle bind and perhaps activate the N2 bond by the same σ-donation and π-back-donation mechanism. The backbonding would be less efficient as compared to low-spin ferrous iron, but in favor of the high-spin case, it would lead to radical character at the ligand, thus perhaps hinting to a mechanism of N2 activation by FeMoCo with transient radical species.

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REFERENCES

(1) CRC Handbook of Chemistry and Physics; Taylor & Francis: Boca Raton, FL, 2012. (2) Einsle, O. Nitrogenase Femo Cofactor: An Atomic Structure in Three Simple Steps. JBIC, J. Biol. Inorg. Chem. 2014, 19, 737−745. (3) Jia, H. P.; Quadrelli, E. A. Mechanistic Aspects of Dinitrogen Cleavage and Hydrogenation to Produce Ammonia in Catalysis and Organometallic Chemistry: Relevance of Metal Hydride Bonds and Dihydrogen. Chem. Soc. Rev. 2014, 43, 547−564. (4) Hu, Y. L.; Ribbe, M. W. Biosynthesis of Nitrogenase Femoco. Coord. Chem. Rev. 2011, 255, 1218−1224. (5) Hoffman, B. M.; Lukoyanov, D.; Yang, Z. Y.; Dean, D. R.; Seefeldt, L. C. Mechanism of Nitrogen Fixation by Nitrogenase: The Next Stage. Chem. Rev. 2014, 114, 4041−4062. (6) Lancaster, K. M.; Roemelt, M.; Ettenhuber, P.; Hu, Y. L.; 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. (7) Thorneley, R. N. F.; Lowe, D. J. Nitrogenase: Substrate Binding and Activation. JBIC, J. Biol. Inorg. Chem. 1996, 1, 576−580. (8) Lowe, D. J.; Thorneley, R. N. F. The Mechanism of KlebsiellaPneumoniae Nitrogenase Action - Pre-Steady-State Kinetics of H-2 Formation. Biochem. J. 1984, 224, 877−886. (9) Lowe, D. J.; Thorneley, R. N. F. The Mechanism of KlebsiellaPneumoniae Nitrogenase Action - the Determination of Rate Constants Required for the Simulation of the Kinetics of N-2 Reduction and H-2 Evolution. Biochem. J. 1984, 224, 895−901. (10) Thorneley, R. N. F.; Lowe, D. J. The Mechanism of KlebsiellaPneumoniae Nitrogenase Action - Pre-Steady-State Kinetics of an Enzyme-Bound Intermediate in N-2 Reduction and of Nh3 Formation. Biochem. J. 1984, 224, 887−894. (11) Thorneley, R. N. F.; Lowe, D. J. The Mechanism of KlebsiellaPneumoniae Nitrogenase Action - Simulation of the Dependences of H-2-Evolution Rate on Component-Protein Concentration and Ratio and Sodium Dithionite Concentration. Biochem. J. 1984, 224, 903− 909. (12) Thorneley, R. N. F.; Lowe, D. J. Nitrogenase of KlebsiellaPneumoniae - Kinetics of the Dissociation of Oxidized Iron Protein from Molybdenum Iron Protein - Identification of the Rate-Limiting Step for Substrate Reduction. Biochem. J. 1983, 215, 393−403. (13) Kirn, J. S.; Rees, D. C. Crystallographic Structure and Functional Implications of the Nitrogenase Molybdenum Iron Protein from Azotobacter-Vinelandii. Nature 1992, 360, 553−560. (14) Lukoyanov, D.; Yang, Z. Y.; Khadka, N.; Dean, D. R.; Seefeldt, L. C.; Hoffman, B. M. Identification of a Key Catalytic Intermediate Demonstrates That Nitrogenase Is Activated by the Reversible Exchange of N-2 for H-2. J. Am. Chem. Soc. 2015, 137, 3610−3615. (15) Lukoyanov, D.; Khadka, N.; Yang, Z. Y.; Dean, D. R.; Seefeldt, L. C.; Hoffman, B. M. Reductive Elimination of H-2 Activates Nitrogenase to Reduce the N N Triple Bond: Characterization of the E-4(4h) Janus Intermediate in Wild-Type Enzyme. J. Am. Chem. Soc. 2016, 138, 10674−10683. (16) Giddey, S.; Badwal, S. P. S.; Kulkarni, A. Review of Electrochemical Ammonia Production Technologies and Materials. Int. J. Hydrogen Energy 2013, 38, 14576−14594. (17) Sellmann, D.; Kunstmann, H.; Knoch, F.; Moll, M. TransitionMetal Complexes with Sulfur Ligands 0.37. Iron, Molybdenum and Ruthenium Complexes with Pentadentate Os4 and Nhs4 Ligands Combining Thiolato, Thioether, and Ether or Amine Donor Functions - Synthesis, Structures, and Reactivity of Co, No, Pme3, and N2h4 Derivatives. Inorg. Chem. 1988, 27, 4183−4190. (18) Li, Y.; Li, Y.; Wang, B. M.; Luo, Y.; Yang, D. W.; Tong, P.; Zhao, J. F.; Luo, L.; Zhou, Y. H.; Chen, S.; Cheng, F.; Qu, J. P. Ammonia Formation by a Thiolate-Bridged Diiron Amide Complex as a Nitrogenase Mimic. Nat. Chem. 2013, 5, 320−326. (19) Yang, D. W.; Li, Y.; Wang, B. M.; Zhao, X. Y.; Su, L. N.; Chen, S.; Tong, P.; Luo, Y.; Qu, J. P. Synthesis and Electrocatalytic Property of Diiron Hydride Complexes Derived from a Thiolate-Bridged Diiron Complex. Inorg. Chem. 2015, 54, 10243−10249.

ASSOCIATED CONTENT

* Supporting Information S

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.9b00177. Mössbauer spectra of all investigated complexes, spin density distribution of 1+, normalized Fe XES Kβ spectra for 1, 1+, 2, and 2+ and Cartesian coordinates [Å] of the optimized model structures for all complexes (PDF) Accession Codes

CCDC 1874850 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 Cambridge Crystallographic Data Centre, 12 Union Road, Cambridge CB2 1EZ, UK; fax: +44 1223 336033.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. ORCID

Thomas Weyhermüller: 0000-0002-0399-7999 Eckhard Bill: 0000-0001-9138-3964 Frank Neese: 0000-0003-4691-0547 Serena DeBeer: 0000-0002-5196-3400 Maurice van Gastel: 0000-0002-1547-6365 Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The Max Planck Society, the European Research Council under the European Union’s Seventh Framework Programme (FP/2007-2013) ERC Grant Agreement No. 615414 (S.D.), and the DFG projects DE 1877/1-1 (S.D.) and NE 690/16-1 (F.N) within the SPP 1927 “Iron−Sulfur for Life” are acknowledged for funding. N


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Inorganic Chemistry (20) Chen, S.; Li, Y.; Yang, D. W.; Luo, L.; Qu, J. P.; Luo, Y. Studies on the Electronic Structure of Thiolate-Bridged Diiron Complexes and Their Single-Electron Reduction Reactions. Chem. Phys. Lett. 2016, 660, 117−122. (21) Anklamm, L.; Schlesiger, C.; Malzer, W.; Grotzsch, D.; Neitzel, M.; Kanngiesser, B. A Novel Von Hamos Spectrometer for Efficient X-Ray Emission Spectroscopy in the Laboratory. Rev. Sci. Instrum. 2014, 85, 053110. (22) Malzer, W.; Grotzsch, D.; Gnewkow, R.; Schlesiger, C.; Kowalewski, F.; Van Kuiken, B.; DeBeer, S.; Kanngiesser, B. A Laboratory Spectrometer for High Throughput X-Ray Emission Spectroscopy in Catalysis Research. Rev. Sci. Instrum. 2018, 89, 113111. (23) Otendal, M.; Tuohimaa, T.; Vogt, U.; Hertz, H. M. A 9keV Electron-Impact Liquid-Gallium-Jet X-Ray Source. Rev. Sci. Instrum. 2008, 79, 016102. (24) Hoszowska, J.; Dousse, J. C.; Kern, J.; Rheme, C. HighResolution Von Hamos Crystal X-Ray Spectrometer. Nucl. Instrum. Methods Phys. Res., Sect. A 1996, 376, 129−138. (25) Neese, F. Software Update: The Orca Program System, Version 4.0. WIREs Comput. Mol. Sci. 2018, 8, e1327. (26) Becke, A. D. Density-Functional Exchange-Energy Approximation with Correct Asymptotic-Behavior. Phys. Rev. A: At., Mol., Opt. Phys. 1988, 38, 3098−3100. (27) van Lenthe, E.; Snijders, J. G.; Baerends, E. J. The Zero-Order Regular Approximation for Relativistic Effects: The Effect of SpinOrbit Coupling in Closed Shell Molecules. J. Chem. Phys. 1996, 105, 6505−6516. (28) van Lenthe, E.; van der Avoird, A.; Wormer, P. E. S. Density Functional Calculations of Molecular Hyperfine Interactions in the Zero Order Regular Approximation for Relativistic Effects. J. Chem. Phys. 1998, 108, 4783−4796. (29) Pantazis, D. A.; Chen, X.-Y.; Landis, C. R.; Neese, F. AllElectron Scalar Relativistic Basis Sets for Third-Row Transition Metal Atoms. J. Chem. Theory Comput. 2008, 4, 908−919. (30) Baerends, E. J.; Ellis, D. E.; Ros, P. Self-Consistent Molecular Hartree-Fock-Slater Calculations - I. The Computational Procedure. Chem. Phys. 1973, 2, 41−51. (31) Römelt, M.; Ye, S. F.; Neese, F. Calibration of Modern Density Functional Theory Methods for the Prediction of Fe-57 Mossbauer Isomer Shifts: Meta-Gga and Double-Hybrid Functionals. Inorg. Chem. 2009, 48, 784−785. (32) Neese, F. Prediction of Electron Paramagnetic Resonance G Values Using Coupled Perturbed Hartree-Fock and Kohn-Sham Theory. J. Chem. Phys. 2001, 115, 11080−11096. (33) Lee, N.; Petrenko, T.; Bergmann, U.; Neese, F.; DeBeer, S. Probing Valence Orbital Composition with Iron K Beta X-Ray Emission Spectroscopy. J. Am. Chem. Soc. 2010, 132, 9715−9727. (34) Pollock, C. J.; Delgado-Jaime, M. U.; Atanasov, M.; Neese, F.; DeBeer, S. K Beta Mainline X-Ray Emission Spectroscopy as an Experimental Probe of Metal-Ligand Covalency. J. Am. Chem. Soc. 2014, 136, 9453−9463. (35) Tsutsumi, K.; Nakamori, H.; Ichikawa, K. X-Ray Mn Kbeta Emission-Spectra of Manganese Oxides and Manganates. Phys. Rev. B 1976, 13, 929−933. (36) Kowalska, J. K.; Hahn, A. W.; Albers, A.; Schiewer, C. E.; Bjornsson, R.; Lima, F. A.; Meyer, F.; DeBeer, S. X-Ray Absorption and Emission Spectroscopic Studies of L2fe2s2 (N) Model Complexes: Implications for the Experimental Evaluation of Redox States in Iron-Sulfur Clusters. Inorg. Chem. 2016, 55, 4485−4497. (37) Pollock, C. J.; Lancaster, K. M.; Finkelstein, K. D.; DeBeer, S. Study of Iron Dimers Reveals Angular Dependence of Valence-toCore X-Ray Emission Spectra. Inorg. Chem. 2014, 53, 10378−10385. (38) Flanagan, H. L.; Singel, D. J. Analysis of N-14 Eseem Patterns of Randomly Oriented Solids. J. Chem. Phys. 1987, 87, 5606−5616. (39) Hinz, A.; Wells, J. S.; Maki, A. G. Heterodyne Frequency Measurements on the Nitric-Oxide Fundamental-Band. J. Mol. Spectrosc. 1986, 119, 120−125.

(40) Karlin, K. D.; Lewis, D. L.; Rabinowitz, H. N.; Lippard, S. J. Addition of Nitrosonium Ion to Mercaptide-Bridged Binuclear Iron(Ii) and Cobalt(Ii) Complexes. J. Am. Chem. Soc. 1974, 96, 6519−6521. (41) Ray, K.; Bill, E.; Weyhermuller, T.; Wieghardt, K. RedoxNoninnocence of the S,S′-Coordinated Ligands in Bis(Benzene-1,2Dithiolato)Iron Complexes. J. Am. Chem. Soc. 2005, 127, 5641−5654. (42) Wallasch, M.; Wolmershauser, G.; Sitzmann, H. Phenolate Complexes of Iron(Ii) in Different Spin States. Angew. Chem., Int. Ed. 2005, 44, 2597−2599.

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Spectroscopic and Quantum Chemical Investigation of Benzene-1,2

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