Valence-Delocalized Diiron(II,III) Cores Supported by Carboxylate-Only Bridging Ligands

In the present communication we disclose the synthesis and characterization of a series of valence-delocalized diiron(II,III) complexes supported only by bridging carboxylate ligands. The existence of this class of compounds first emerged during an investigation of the dioxygen reactivity of tetra(carboxylate)diiron(II) complexes as models for dioxygen-activating centers in non-heme diiron enzymes.1,2 An intriguing green S ) /2 paramagnetic component was detected in the metastable mixture formed in the reaction of a diiron(II) complex with dioxygen, preliminary Mossbauer studies of which indicated it to be a valence-delocalized diiron(II,III) cluster.2 An independent synthesis of this green species and related complexes has now been achieved, spin-coupling in which apparently is dominated by a double exchange mechanism.3,4 These findings provide the first observation of this phenomenon in diiron clusters having no single-atom bridging ligand(s), and suggest that a direct metalmetal interaction promoted by the short intermetal distance may help achieve the spin delocalization. The diiron(II) precursor compound [Fe2(μ-O2CAr)4(4-BuC5H4N)2] (1)2 undergoes a reversible one-electron oxidation (E1/2 ) -216 mV vs Cp2Fe/Cp2Fe; ∆Ep ) 89 mV, scan rate ) 25 mV/s) as revealed by cyclic voltammetry in CH2Cl2 solution (Figure S1). Compound 1 can be oxidized chemically, either by [Cp2Fe](PF6) or by AgOTf, in CH2Cl2 (Scheme 1) to generate dark emerald green solutions of [Fe2(μ-O2CAr)4(4-BuC5H4N)2]X (X ) PF6 or OTf-) (4). Compound 4 is thermally stable but decolorizes upon exposure to dioxygen or addition of coordinating solvents (THF, MeCN, or MeOH). The analogous green mixedvalence diiron compounds [Fe2(μ-O2CAr)4(py)2](OTf) (5) and [Fe2(μ-O2CAr)4(THF)2](PF6) (6) were prepared from the corresponding diiron(II) precursors, [Fe2(μ-O2CAr)2(O2CAr)2(py)2] (2) and [Fe2(μ-O2CAr)2(O2CAr)2(THF)2] (3),1 by the same route (Scheme 1). The structure of 4 (Figure S2), as determined by X-ray crystallography,5 reveals a shortening by ∼0.11 A of the Fe‚‚‚Fe separation (2.713(3) A) compared with that in the precursor compound 1 (2.823(1) A).2 Each iron center in 4 retains square-pyramidal coordination with negligible geometric changes upon 1-electron oxidation of the diiron(II) complex, consistent with the reversible nature of the redox process as observed by cyclic voltammetry. In contrast, the pyridine (2) and THF (3) derivatives undergo significant structural reorganization upon 1-electron oxidation. In both cases, the two carboxylate ligands in the precursor complex shift from terminal chelating to bidentate bridging (Scheme 1), with concomitant shortening of the Fe‚‚‚Fe separations from 4.2189(13) A to 2.6982(13) A (5) and from 4.2822(7) A to 2.6633(11) A (6) (Figures S3 and S4). As illustrated in Figure 1, the Fe-N/Fe-O distances are similar for both iron centers in 5 at 188 K (average Fe1-O, 2.011(7) A; Fe2-O, 2.002(6) A), consistent with the valence-delocalized nature of the diiron cluster. The electronic spectrum of 4 in CH2Cl2 (Figure S5) displays a broad intervalence charge-transfer band at 670 nm ( ≈ 3200 M-1 cm-1). Similar broad visible absorptions were observed for 5 and 6. The values of ∆ν1/2 calculated by Hush’s relationship6,7 for class II mixed-valence compound are much larger than those measured for 4-6 (Table S1), indicating the valence-delocalized nature of these centers. The S ) /2 ground-state resulting from the high-spin iron(II)iron(III) unit is characterized by a broad g ) 10 signal in the EPR spectrum displayed by a frozen CH2Cl2 solution of 4.8 Consistent with this result, the μeff value of 11.0 μB (300 K) obtained from the magnetic susceptibility data (Figure S6) on a powder sample of 4 is close to the spin-only value of 9.9 μB (g ) 2.00) for the S ) /2 total spin system. Similar values were obtained for 5 and 6 (Table S1).9 The electronic structure of 4 was further probed by Mossbauer studies. At 4.2 K and in a weak applied magnetic field (50 mT), † Massachusetts Institute of Technology. ‡ Emory University. § Carnegie Mellon University. (1) Lee, D.; Lippard, S. J. J. Am. Chem. Soc. 1998, 120, 12153-12154. (2) Lee, D.; Du Bois, J.; Petasis, D.; Hendrich, M. P.; Krebs, C.; Huynh, B. H.; Lippard, S. J. J. Am. Chem. Soc. 1999, 121, 9893-9894. (3) Munck, E.; Papaefthymiou, V.; Surerus, K. K.; Girerd, J.-J. Metal Clusters in Proteins; Que, L., Jr., Ed.; American Chemical Society: Washington, DC, 1988; pp 302-325. (4) Blondin, B.; Girerd, J.-J. Chem. ReV. 1990, 90, 1359-1376. (5) Crystal data: triclinic, P1h, a ) 16.472(5) A, b ) 16.706(5) A, c ) 17.888(5) A, R ) 108.243(5)°, â ) 90.614(5)°, γ ) 90.587(5)°, Z ) 2, T ) -85 °C, R ) 0.104. This structure was not fully refined due to severe disorder in the lattice solvent molecules and the PF6 counterion. (6) Hush, N. S. Prog. Inorg. Chem. 1967, 8, 391-444. (7) Creutz, C. Prog. Inorg. Chem. 1983, 30, 1-73. (8) This signal is identical with that recorded for the S ) /2 component in the paramagnetic species generated from the reaction of 1 with dioxygen at low temperature, detailed parameters of which were reported previously.2 (9) Detailed analysis of the magnetic susceptibility data is currently in progress. Figure 1. ORTEP diagram of the [Fe2(μ-O2CAr)4(py)2](OTf) (5) cation showing 50% probability thermal ellipsoids for non-hydrogen atoms. For clarity, all atoms of the 2,6-di(p-tolyl)benzoate ligands, except for the carboxylate groups and the R-carbon atoms, were omitted. Selected interatomic distances (A): Fe1‚‚‚Fe2, 2.6982(13); Fe1-O1A, 2.011(3); Fe1-O1B, 2.008(3); Fe1-O1C, 2.013(4); Fe1-O1D, 2.012(4); Fe1N1, 2.104(4); Fe2-O2A, 2.015(3); Fe2-O2B, 1.989(3); Fe2-O2C, 2.000(3); Fe2-O2D, 2.005(3); Fe2-N2, 2.085(4).