Synthesis, Properties, and Reactivity of Diketiminate-Supported Cobalt

Nov 13, 2009 - Synopsis. A three-coordinate cobalt(II) precursor leads to unusual cobalt(II) fluoride complexes, including an interesting trimetallic ...
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Organometallics 2009, 28, 6650–6656 DOI: 10.1021/om900394t

Synthesis, Properties, and Reactivity of Diketiminate-Supported Cobalt Fluoride Complexes Keying Ding, Thomas R. Dugan, William W. Brennessel, Eckhard Bill,* and Patrick L. Holland* Department of Chemistry, University of Rochester, Rochester, New York 14627, and Max-Planck-Institut f€ ur Bioanorganische Chemie, M€ ulheim an der Ruhr D45470, Germany Received May 14, 2009

Reaction of the three-coordinate cobalt(II) methyl complex LtBuCoCH3 (LtBu = bulky β-diketiminate ligand) with (CH3)3SnF gives an interesting trimetallic complex, [LtBuCo(μ-F)2]2Sn(CH3)2 (1). Magnetic susceptibility studies indicate that 1 consists of two mixed-valence cobalt(II/III) fluoride units bridged by the organometallic tin(III) fragment [SnF2(CH3)2]-. The addition of pyridine releases the tin bridge, and the cobalt fluoride species can be isolated as the high-spin cobalt(II) complex LtBuCoF(py) (2). Treating 2 with triethylsilane in an effort to form LtBuCoH(py) gave instead the high-spin cobalt(I) complex LtBuCo(py), with elimination of H2. Attempts at catalytic or stoichiometric hydrodefluorination of fluorinated aromatics with silane using 2 as a catalyst were unsuccessful. The lack of hydrodefluorination activity may be due to H2 elimination from lowcoordinate cobalt hydride complexes or from inhibition by a fourth donor to the metal.

Introduction Organometallic fluoride complexes have received considerable attention due to their interesting structures and reactions.1 Organometallic complexes containing a fluorine ligand are of special interest since fluoride-containing complexes are catalytic intermediates in C-F bond activation reactions. The reduction of fluorocarbons with transition metal catalysts is an active research area with many recent advances.2 Inexpensive first-row late transition metals are excellent candidates for new catalysts because of their abundance and low cost.3,4 In this context, we have reported low-coordinate iron(II) fluoride complexes that come from *Corresponding authors. E-mail: [email protected]. (1) (a) Murphy, E. F.; Murugavel, R.; Roesky, H. W. Chem. Rev. 1997, 97, 3425–3468. (b) Roesky, H. W. Inorg. Chem. 1999, 38, 5934–5943. (c) Grushin, V. V. Chem. Eur. J. 2002, 8, 1006–1014. (2) (a) Kiplinger, J. L.; Richmond, T. G.; Osterberg, C. E. Chem. Rev. 1994, 94, 373–431. (b) Douvris, C.; Ozerov, O. V. Science 2008, 321, 1188– 1190. (c) Reade, S. P.; Mahon, M. F.; Whittlesey, M. K. J. Am. Chem. Soc. 2009, 131, 1847–1861. (d) Meier, G.; Braun, T. Angew. Chem., Int. Ed. 2009, 48, 1546–1548. (e) Amii, H.; Uneyama, K. Chem. Rev. 2009, 109, 2119–2183. (3) Representative recent examples of C-F activation by Fe, Co, and Ni complexes: (a) Schaub, T.; Radius, U. Chem. Eur. J. 2005, 11, 5024– 5030. (b) Li, X.; Sun, H.; Yu, F.; Floerke, U.; Klein, H.-F. Organometallics 2006, 25, 4695–4697. (c) Schaub, T.; Backes, M.; Radius, U. J. Am. Chem. Soc. 2006, 128, 15964–15965. (d) Schaub, T.; Backes, M.; Radius, U. Eur. J. Inorg. Chem. 2008, 2680–2690. (e) Schaub, T.; Fischer, P.; Steffen, A.; Braun, T.; Radius, U.; Mix, A. J. Am. Chem. Soc. 2008, 130, 9304–9317. (f) Nakao, Y.; Kashihara, N.; Kanyiva, K. S.; Hiyama, T. J. Am. Chem. Soc. 2008, 130, 16170–16171. (g) Johnson, S. A.; Huff, C. W.; Mustafa, F.; Saliba, M. J. Am. Chem. Soc. 2008, 130, 17278–17280. (4) Vela, J.; Smith, J. M.; Yu, Y.; Ketterer, N. A.; Flaschenriem, C. J.; Lachicotte, R. J.; Holland, P. L. J. Am. Chem. Soc. 2005, 127, 7857– 7870. (5) (a) Herzog, A.; Liu, F.-Q.; Roesky, H. W.; Demsar, A.; Keller, K.; Noltemeyer, M.; Pauer, F. Organometallics 1994, 13, 1251–1256. (b) Dorn, H.; Roesky, H. W.; Morris, R. J. Inorg. Synth. 2002, 33, 234–239. pubs.acs.org/Organometallics

Published on Web 11/13/2009

reaction of an iron(II) alkyl complex with the interesting fluoride source (CH3)3SnF.1b,4,5 These low-coordinate iron complexes catalyze hydrodefluorination of fluorinated aromatic and olefinic compounds using Et3SiH (Scheme 1). The hydrodefluorination reaction proceeds to completion with 20% loading of LtBuFeF (where LtBu represents the β-diketiminate ligand 2,2,6,6-tetramethyl-3,5-bis(2,6-diisopropylphenylimido)hept-4-yl). Despite the mild conditions of the iron-catalyzed reaction, it never progressed beyond 5-10 turnovers, suggesting that the catalyst is readily deactivated. As one potential way to modify the complexes and increase their stability, we chose to explore the use of the analogous cobalt complexes of LtBu. We present here the results of these studies, which have yielded some unusual and interesting cobalt(II) fluoride complexes.

Results and Discussion Synthesis and Properties of a Cobalt Tin Fluoride Complex. As previously reported, treatment of LtBuCoCl with methylmagnesium chloride gives the three-coordinate cobalt(II) methyl complex LtBuCoCH3.6 This methyl complex has been crystallographically characterized.6 Its solution magnetic moment (4.9 μB) and proton NMR spectra are similar to those of LtBuCoCl, showing that it also has a high-spin d7 configuration at cobalt. We envisioned that LtBuCoCH3 could be a precursor to a cobalt fluoride complex, by analogy with the reaction of LtBuFeCH3 with (CH3)3SnF to give LtBuFeF.4 Reaction of LtBuCoCH3 with 1 molar equiv of (CH3)3SnF at a range (6) Holland, P. L.; Cundari, T. R.; Perez, L. L.; Eckert, N. A.; Lachicotte, R. J. J. Am. Chem. Soc. 2002, 124, 14416–14424. r 2009 American Chemical Society

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Scheme 1. Catalytic C-F Activation by the Low-Coordinate Iron Fluoride Complex LtBuFeF4 a

a

Ar represents a 2,6-diisopropylphenyl substituent.

Scheme 2. Synthesis of 1a Figure 1. ORTEP diagram of [LtBuCo(μ-F)2]2Sn(CH3)2 (1) using 50% thermal ellipsoids. a

Ar represents a 2,6-diisopropylphenyl substituent.

of temperatures between 25 and 100 C and in different solvents (toluene, diethyl ether) gives a mixture as judged by 1 H NMR spectroscopy. There is partial conversion to a new compound, and unreacted LtBuCoCH3 is present some amount of unreacted LtBuCoCH3. The yield of the new compound increases when the ratio of (CH3)3SnF to LtBuCoCH3 is increased. All of the LtBuCoCH3 is consumed when it is stirred with 2 equiv of (CH3)3SnF in toluene at 80 C for 24 h. Under these conditions, only one paramagnetic product is observed by 1H NMR spectroscopy. (CH3)3SnF does not react with the product; therefore, slightly more than 2 equiv of (CH3)3SnF are typically used to ensure that all LtBuCoCH3 is converted to product. Unreacted (CH3)3SnF (which is insoluble in toluene) is easily removed from the reaction mixture by filtration. The cobalt-containing product (1) is isolated in 65% yield as a bright red powder that readily dissolves in solvents such as diethyl ether, toluene, or THF. X-ray diffraction analysis (see below) identifies it as the trimetallic complex [LtBuCo(μ-F)2]2Sn(CH3)2 3 2Et2O (1 3 2Et2O). The isolation of this Co2Sn product in high yield implies that the stoichiometry of the reaction is as shown in Scheme 2. The volatile byproduct Sn(CH3)4 is removed during evaporation of solvent. The (undoubtedly complicated) mechanism of the reaction is unknown at this time; no intermediate species are seen by 1H NMR spectroscopy during the course of the reaction. The diethyl ether solvate 1 3 2Et2O crystallized in space group P1 with Z = 2. There is no crystallographic symmetry within the molecule. The structure features two cobalt atoms and one tin atom held together by fluoride bridges (Figure 1). The tin atom is six-coordinate by virtue of two methyl groups in addition to the four bridging fluorides. Each cobalt atom has a four-coordinate tetrahedral geometry and is bound by a β-diketiminate ligand. The [Co(μ-F)2Sn] rhombs in 1 are distinctly asymmetric. In each rhomb, one of the Co-F distances is shorter (Co1-F1 1.9340(9) A˚, Co2-F3 1.9315 A˚), and the other is longer (Co1-F2 1.9941(9) A˚, Co2-F4 1.9985(9) A˚).7 In each case, the fluoride with the longer bond to cobalt (7) The average length of bridging Co-F bonds in the Cambridge Structural Database is 2.06(4) A˚. Cambridge Structural Database, ConQuest version 1.7, November 2008 release. Allen, F. H. Acta Crystallogr. 2002, B58, 380–388.

Figure 2. Side view of [LtBuCo(μ-F)2]2Sn(CH3)2 (1), showing the asymmetry of the bonding around the tin atom (which lies in the middle, hidden behind C2). The substituents on the diketiminate ligands have been removed for clarity.

forms a shorter bond to the tin atom (Sn1-F2 2.0964(9) A˚, Sn1-F4 2.1022(9) A˚) than does the other fluoride (Sn1-F1 2.2546(9) A˚, Sn1-F3 2.2949(9) A˚). Therefore, the structure suggests a strongly covalent SnF2(CH3)2 species (containing Sn1, C1, C2, F2, and F4) that bridges two cobalt-fluoride units. The tin atom has a distorted octahedral geometry with the two methyl groups occupying trans positions. The Sn-C bond lengths of 2.088(2) and 2.091(2) A˚ are nearly identical to the Sn-C bond lengths in (NH4)2[(CH3)2SnF4] and K2[(CH3)2SnF4].8 The two short Sn-F bonds are cis, breaking the inversion symmetry of the molecule. Magnetism of Complex 1. This complex is paramagnetic at room temperature. The paramagnetism of 1 leads to highly shifted resonances between δ 64 and -53 ppm in its 1H NMR spectrum. These shifts are similar to those in other paramagnetic cobalt complexes of LtBu.6 Eight resonances are observed in the 1H NMR spectrum, which integrate as expected for D2d symmetry. A resonance at δ -38 ppm integrating to six protons is assigned to the tin-bound methyl groups. The presence of seven additional resonances shows that the β-diketiminate ligands are equivalent and have local C2v symmetry (giving seven inequivalent proton environments). Therefore, the asymmetry of Co-F and Sn-F bond lengths observed in the structure is averaged on the 1H NMR time scale to make all Co-F bonds equivalent. The solid-state magnetization of 1 from a SQUID magnetometry experiment is shown in Figure 3 (top trace, black circles). The effective magnetic moment shows a faint maximum at about 40 K and drops rapidly below 30 K because of field saturation and zero-field splitting, but passes over to a very shallow minimum at higher temperatures. The appearance of this curve seems at first glance to be consistent with a (8) Ahmed, I. A.; Kastner, G.; Reuter, H.; Schultze, D. J. Organomet. Chem. 2002, 649, 147–151.

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Table 1. Selected Interatomic Distances and Angles in Diketiminate Cobalt(II) Complexes Co-F (A˚) Co-N(dik) (A˚) N(dik)-Co-N(dik) (deg) Co-N(py) (A˚) N(dik)-Co-N(py) (deg) Sn-F (A˚) Sn-C (A˚)

[LtBuCo(μ-F)2]2Sn(CH3)2 (1)

LtBuCoF(py) (2)

1.9340(9), 1.9315(9), 1.9941(9), 1.9985(9) 1.943(1), 1.943(1) 99.82(5), 98.95(5)

1.843(2) 1.969(3), 1.962(3) 97.35(11) 2.052(3) 119.3(1), 115.1(1)

LtBuCo(py) (3)

1.902(1), 1.908(1) 98.92(5) 1.968(1) 129.24(5), 130.18(5)

2.0964(9), 2.1022(9), 2.2546(9), 2.2949(9) 2.088(2), 2.091(2)

Figure 3. Temperature dependence of the effective magnetic moment of 1 recorded at a field of B = 1 T (top trace, black circles, χdia and χTIP = 90  10-6 emu subtracted). The red line is the preferred fit with three spins as indicated with parameters JCo-Sn = -24 cm-1, JCo-Co = þ8.8 cm-1, gSn = 2.23, gCo = 2.106, and DCo ≈ -12.1 cm-1. The lower trace (gray circles) shows the same data corrected for a TIP-like temperatureindependent contribution by subtracting χTIP = 4533  10-6 emu, such that a fit to a two-spin model is possible with SCo = 3/2 coupled with J = þ20 cm-1, gCo = 2.016, DCo = þ10 cm-1 (blue line). The lower two-spin model is untenable (see text).

weak ferromagnetic interaction between the two cobalt ions, but the high-temperature values of ca. 7 μB are not consistent with the presence of two spins of SCo = 3/2 for Co(II), but seem to indicate Co(III) with SCo = 2 ðμeff ¼ p ffiffiffi rather pffiffiffiffiffiffiffiffiffiffiffi 2 3 g 3 2  3≈6:93 μB Þ for g = 2 in the high-temperature limit of virtually uncoupled spins). However, the position of the maximum, which is a measure of the coupling strength,9 and its height and overall shape of the curve cannot be reasonably simulated with any choice of a coupling constant J and zero-field splitting parameter DCo for two cobalt(II) or two cobalt(III) ions. A mediocre fit of the data with two spins SCo = 3/2 is possible only if the data are corrected by subtraction of a tremendous contribution from temperature-independent paramagnetism (TIP) of χTIP ≈ 4300  10-6 emu (Figure 3, lower trace in gray and simulation in blue, JCo-Co = þ20 cm-1, DCo = þ10 cm-1, gCo = 2.016). Such a large TIP contribution is unreasonable.10 Moreover, it would be difficult to imagine why the cobalt ions in 1 would be ferromagnetically coupled, because the orthogonality of magnetic orbitals that is usually necessary to suppress the common antiferromagnetic interaction is not plausible here because of the tetrahedral geometry of the cobalt ions. (9) Kahn, O. Molecular Magnetism; VCH: Weinheim, Germany, 1993. (10) Such a TIP contribution is too large, even if one invokes a major contamination of the sample with superparamagnetic cobalt oxide particles with single-domain behavior that leads to temperature-independent magnetization.

Another possibility is that tin is reduced to Sn(III) due to an internal redox process that oxidizes one cobalt to Co(III). In this model, compound 1 is a mixed-valence trimer with spins S1 = 1/2 at Sn(III), S2 = 3/2 at Co(II), and S3 = 2 at Co(III). A corresponding simulation yields an excellent fit (Figure 3, top trace in red). The ground state, resulting from the optimized parameters JSn-Co = -24 cm-1, JCo ≈ þ9 cm-1, and DCo ≈ -12 cm-1, has a total spin Stotal = 3 from ferromagnetic coupling of the two Co spins, which are both antiferromagnetically coupled to Sn(III) (Stotal = 3/2 þ 2 1/2 = 3). In the above fit, the two J values for the Sn-Co coupling and the two D values for the cobalt ions are constrained to be equal. This is done not only to reduce the number of free parameters in the fit but also because the Co(II)/Co(III) sites are likely to be valence-delocalized. Two pieces of evidence support the idea of valence delocalization between the two cobalt ions. First, the two cobalt sites are indistinguishable in the solid-state molecular structure. Second, the fit to the magnetic susceptibility data yields a positive value for JCo-Co, the spin coupling of the two cobalt sites; a ferromagnetic interaction of this sort is commonly observed in high-spin systems with valence delocalization.11 Although tin(III) is rare in isolable monometallic species,12 a related system [LFeSnFeL]3þ (L = 1,4,7-tris(4-tert-butyl-2-mercaptobenzyl)-1,4,7-triazacyclononane) with thiolate bridges has been examined with complementary 57Fe- and 119SnM€ ossbauer spectroscopy, which clearly showed that an internal redox shift gives a octahedral bridging Sn(III) ion between a pair of coupled low-spin Fe(II)/Fe(III) ions.13 This close analogy supports the model of two cobalt(II/III) ions bridged by tin(III) that most clearly explains the magnetic data. Monomeric Four-Coordinate Cobalt(II) Fluoride Complex. Treating 1 with 2 equiv of pyridine gives an immediate color change from red to yellow and the formation of a white precipitate. Light yellow crystals of LtBuCoF(py) (2) could be isolated from the filtrate in 72% yield. The overall reaction is shown in Scheme 3. The loss of SnF2(CH3)2 from 1 is consistent with the long bonds found between the central SnF2(CH3)2 unit and the LtBuCoF units in the X-ray crystal structure of 1. Though the magnetism of 1 indicates an (11) (a) Bominaar, E. L.; Hu, Z.; M€ unck, E.; Girerd, J.-J.; Borshch, S. J. Am. Chem. Soc. 1995, 117, 6976–6989. (b) Borshch, S. A.; Bominaar, E. L.; Blondin, G.; Girerd, J.-J. J. Am. Chem. Soc. 1993, 115, 5155–5168. (c) Bominaar, E. L.; Borshch, S. A.; Girerd, J.-J. J. Am. Chem. Soc. 1994, 116, 5362–5372. (12) (a) Davidson, P. J.; Hudson, A.; Lappert, M. F.; Lednor, P. W. J. Chem. Soc., Chem. Commun. 1973, 829–30. (b) Power, P. P. Chem. Rev. 2003, 103, 789–809. (c) Becker, M.; Foerster, C.; Franzen, C.; Hartrath, J.; Kirsten, E.; Knuth, J.; Klinkhammer, K. W.; Sharma, A.; Hinderberger, D. Inorg. Chem. 2008, 47, 9965–9978. (13) Glaser, T.; Bill, E.; Weyherm€ uller, T.; Meyer-Klaucke, W.; Wieghardt, K. Inorg. Chem. 1999, 38, 2632–2642.

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Figure 5. Idealized symmetry and occupation of electrons in a tetrahedral crystal field. Because the idealized tetrahedron gives an E ground state in the Fe2þ compound, it should be subject to a Jahn-Teller distortion (splitting on the right).

Figure 4. ORTEP diagram showing one of the two independent molecules in the crystal structure of LtBuCoF(py) (2), using 50% thermal ellipsoids. Scheme 3. Synthesis of 2a

a

Ar represents a 2,6-diisopropylphenyl substituent.

internal redox reaction, this electronic redistribution is rapidly reversed upon release of the tin(IV) species. This electronic redistribution is analogous to the easy ejection of N2 from LtBuFeNNFeLtBu, though its ground state is best described as two iron(II) ions bridged by the N22- dianion.14 The loss of the insoluble tin byproduct occurs only upon addition of a Lewis base. The need for a Lewis base to eject a tin fluoride contrasts with the iron compound LtBuFeF, which does not react with excess (CH3)3SnF used in its synthesis.4 X-ray crystallography indicates the molecular structure of 2. There are two independent molecules in the asymmetric unit with similar metrical parameters, and one of these is shown in Figure 4. The Co-F distances in the two independent molecules are 1.843(2) and 1.852(2) A˚, and these terminal Co-F bonds are much shorter than the bridging Co-F bonds in 1. The coordination geometry around cobalt is pseudotetrahedral, and the chelating diketiminate ligand restricts one angle (N-Co-N) to 97-98. It is interesting to compare this structure to the related iron(II) fluoride pyridine adduct,4 where there is a pyramidal distortion on the iron coordination geometry. The extent of this pyramidal distortion can be quantified using τ,15 which has ideal values of τ = 1 for perfect trigonal-pyramidal geometry (where the metal sits in the plane of the three basal ligands) and τ = 0 for tetrahedral geometry.16 The τ value for LtBuFeF(4-tBu-py) is (14) (a) Smith, J. M.; Sadique, A. R.; Cundari, T. R.; Rodgers, K. R.; Lukat-Rodgers, G.; Lachicotte, R. J.; Flaschenriem, C. J.; Vela, J.; Holland, P. L. J. Am. Chem. Soc. 2006, 128, 756–769. (b) Stoian, S. A.; Vela, J.; Smith, J. M.; Sadique, A. R.; Holland, P. L.; M€unck, E.; Bominaar, E. L. J. Am. Chem. P P Soc. 2006, 128, 10181–10192. (15) τ = [ (Lbasal-Fe-Lbasal) - (Lbasal-Fe-Laxial)]/90. (16) Vela, J.; Stoian, S.; Flaschenriem, C. J.; M€ unck, E.; Holland, P. L. J. Am. Chem. Soc. 2004, 126, 4522–4523.

0.45, and that of LtBuFeF(NCCH3) is 0.34.4 The τ value for each independent cobalt atom in 2 is less than 0.1, showing that 2 lacks the trigonal-pyramidal distortion found in its iron analogue. It is possible to explain the differences in trigonalpyramidal distortion between the iron and cobalt species in terms of a simple crystal-field model that approximates the metal geometry as tetrahedral (Figure 5). The high-spin d6 configuration of iron(II) places a single electron in the lower e level. Because this set of orbitals would be degenerate, the complex should be subject to a Jahn-Teller distortion to remove the degeneracy and split these orbitals (not shown).17 In the high-spin d7 configuration of cobalt(II), on the other hand, the e pair is completely filled, and no Jahn-Teller distortion is expected. This rationalizes why LtBuCoF(py) lacks the pyramidal distortion previously seen in LtBuFeF(py). Compound 2 has a solution magnetic moment of 4.7(4) μB, consistent with a high-spin CoII center with three unpaired electrons and significant orbital angular momentum. Additional support for an S = 3/2 ground state comes from X-band EPR spectroscopy of a frozen toluene solution at 8 K (Figure 6). The rhombic signal shows components at geff = 6.8, 3.5, and 1.4. The unusually large noise in the spectrum is due to its breadth; the signal spreads over a wide range from 1000 to 7000 G. Reaction of LtBuCoF(py) 3 2Et2O (2) with Et3SiH. Transition metal fluorides may react with Si-H-containing compounds to provide hydride complexes, because of the great thermodynamic driving force to form Si-F bonds.18 This method has been used in the analogous diketiminate iron systems to give iron(II) hydride complexes.4,19 When 2 is heated with triethylsilane for 7 h at 50 C, one major paramagnetic product is observed by 1H NMR spectroscopy. However, instead of the anticipated product LtBuCoH(py), the product is the cobalt(I) complex LtBuCo(py) (3) (Scheme 4). Compound 3 can also be synthesized in 59% yield by adding pyridine to a solution of the recently published Co(II) hydride complex LtBuCo(μ-H)2CoLtBu.20 Although the elemental analysis of crystallized samples of 3 from each route suggest that small impurities are present, the identification of compound 3 is clearly supported by X-ray crystallography. (17) In a previous publication, we noted that the pyramidal distortion in four-coordinate iron(II) complexes is greater than iron(III) complexes, which is also attributable to Jahn-Teller distortions in d6 ions. Vela, J.; Cirera, J.; Smith, J. M.; Lachicotte, R. J.; Flaschenriem, C. J.; Alvarez, S.; Holland, P. L. Inorg. Chem. 2007, 46, 60–71. (18) (a) Murphy, E. F.; Murugavel, R.; Roesky, H. W. Chem. Rev. 1997, 97, 3425–3468. (b) Doherty, N. M.; Hoffman, N. W. Chem. Rev. 1991, 91, 553–573. (c) Hoffman, N. W.; Prokopuk, N.; Robbins, M. J.; Jones, C. M.; Doherty, N. M. Inorg. Chem. 1991, 30, 4177–4181. (d) Doherty, N. M.; Critchlow, S. C. J. Am. Chem. Soc. 1987, 109, 7906–7908. (19) Dugan, T. R.; Holland, P. L. J. Organomet. Chem. 2009, 694, 2825–2830. (20) Ding, K.; Brennessel, W. W.; Holland, P. L. J. Am. Chem. Soc. 2009, 131, 10804–10805.

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Figure 6. X-band EPR spectrum of a frozen solution of 2 in toluene (8 K). The geff values of different components are indicated.

The crystal structure of 3 is shown in Figure 7. In 3, the cobalt atom has a pseudotrigonal-planar geometry, with a sum of angles around cobalt of 358.34(8). The pyridine nitrogen atom is very close to the C2 axis of the diketiminate-

Figure 7. ORTEP diagram of LtBuCo(py) (3), using 50% thermal ellipsoids.

Scheme 4. Synthesis of 3

cobalt unit, with N(diketiminate)-Co-N(pyridine) angles of 129.24(5) and 130.18(5). This kind of Y shape has been seen in three-coordinate nickel(II) diketiminate complexes that have a high-spin d8 configuration.6,21 The Co-N(pyridine) distance, 1.968(1) A˚, is slightly longer than the Co-N(diketiminate) distances of 1.902(1) and 1.908(1) A˚, but is shorter than the Co-N(pyridine) distance in 2 of 2.052(3) A˚. The pyridine ring bends away from the CoN3 plane, giving an angle Co-N14-C34 = 152.35(7) between the Co-N bond and the N-C(para) vector in the pyridine ligand (Figure 8). Considering that hydrides are often difficult to locate in X-ray crystal structures, the potential presence of hydrides has been evaluated using multiple strategies. First, there is no peak in the difference Fourier map of the crystal structure that is assignable to a cobalt hydride ligand. Additionally, the planar disposition of the three cobalt-bound atoms strongly suggests a three-coordinate metal center. In addition, gas chromatography was used to detect 0.65 equiv of H2 produced during the reaction of LtBuCoF(py) and Et3SiH. Therefore, there is strong support for the reductive elimination of hydride ligands to give H2 during the formation of 3. Compound 3 has a solution magnetic moment of 3.5 ( 0.3 μB. This value is consistent with the S = 1 ground state of high-spin cobalt(I). The 1H NMR spectrum has paramagne(21) Melzer, M. M.; Jarchow-Choy, S.; Kogut, E.; Warren, T. H. Inorg. Chem. 2008, 47, 10187–10189.

Figure 8. Side view of LtBuCo(py) (3), showing the planar cobalt atom and the distortion of the pyridine ring out of the CoN3 plane.

tically shifted resonances with integrations that are consistent with C2v symmetry, as expected from the X-ray crystal structure if the pyridine can rotate and/or dissociate reversibly on the 1H NMR time scale. Attempted Hydrodefluorination Reactions. We have shown that low-coordinate iron fluoride complexes are capable of catalytically reducing the C-F bonds of unsaturated fluorocarbons.4 Therefore, we tested cobalt fluoride complexes 1 and 2 for catalytic hydrodefluorination activity. We used 19 F{1H} NMR and 1H NMR spectroscopies to monitor the reactions of triethylsilane (Et3SiH) with hexafluorobenzene (C6F6) or perfluoropyridine in the presence of 0.2 molar equiv of complex 1 or 2.22 Heating at 50 or 80 C for several days in C6D6 showed no significant amounts of reduced fluoroaromatics. In the case of 1, one problem was clearly the stability of the compound. A mixture of 1 and 2 equiv of Et3SiH heated to 80 C decomposed to unidentifiable products within 1 day, whether in the presence or the absence of fluoroaromatic substrates. Why is the C-F activation activity poorer in the LtBuCoF systems than in the LtBuFeF system? While the reason is not certain, two potential explanations can be identified. First, the iron complex is three-coordinate, while isolation of the cobalt complexes requires a fourth donor (F2Sn(CH3)2 in 1 (22) In the absence of a catalyst there was no sign of reaction between any of the fluorinated substrates and trisubstituted silanes.

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and pyridine in 2). The presence of a donor may inhibit the hydrodefluorination reaction. In order to test this idea, the catalytic hydrodefluorination activity of four-coordinate LtBuFeF(py) was examined. Heating Et3SiH and C6F6 at 45 C for 4 days with 0.2 equiv of the pyridine-bound iron complex LtBuFeF(py) does not result in the formation of any C6HF5. This is in contrast to the effectiveness of LtBuFeF as a precatalyst under the same conditions, where C6HF5 is produced with 2.5 turnovers.4 In addition, LtBuFeF(py) does not react stoichiometrically with Et3SiH to give LtBuFeH(py), while LtBuFeF does react with Et3SiH to give [LtBuFeH]2.19 Therefore, pyridine effectively inhibits ironbased HDF catalysis, and it is reasonable to guess that any potential cobalt-based catalysis could also be inhibited by the inevitable donors present in solution. Second, Co-H species in the catalytic reaction may eliminate H2, providing an unproductive pathway for loss of hydride intermediates in the hydrodefluorination reaction. This idea is supported by the rapid elimination of H2 during the formation of LtBuCo(py) (3) described above. In a separate publication, we have shown that N2 is capable of displacing H2 from isolated cobalt(II) and cobalt(I) hydride complexes.20 Therefore, it is possible that any productive reaction of Et3SiH with LtBuCoF species to form a cobalt hydride would be followed rapidly by hydride elimination as H2. This contrasts with the iron hydrides, which do not thermally eliminate H2.23 Therefore, the greater stability of hydride intermediates in the iron system correlates with the ability of the iron complexes to perform catalytic hydrodefluorination.

Conclusions Despite the superficial similarity of the three-coordinate iron(II) and cobalt(II) diketiminate complexes such as LtBuMCl and LtBuMCH3, the reactions of these compounds with tin fluorides give products with distinct differences. A three-coordinate iron(II) fluoride complex LtBuFeF was isolable, but its three-coordinate cobalt(II) fluoride analogue has not been accessible using similar preparatory methods. An interesting disproportionation of Sn-F and Sn-C bonds yields an unusual heterotrimetallic compound in which an SnF2(CH3)2 unit bridges two cobalt-fluoride units. The solid-state magnetic susceptibility of this compound suggests that it has transferred an electron from cobalt(II) to tin(IV) to give an interesting species with tin(III) bridging two valence-delocalized cobalt(II/III) ions. However, this redox shift is reversible, and a tin(IV) species can be displaced by pyridine to give a monomeric four-coordinate cobalt(II) fluoride complex. Both cobalt complexes are incapable of performing hydrodefluorination reactions. Although multiple explanations for this lack of catalytic activity are possible, the most likely reasons for the lower activity in the current system are (1) the instability of cobalt hydrides with respect to loss of H2 through reductive elimination and (2) the presence of a fourth ligand on the cobalt complexes.

Experimental Section General Procedures. All manipulations were performed under a nitrogen or argon atmosphere in an M. Braun glovebox (23) Yu, Y.; Sadique, A. R.; Smith, J. M.; Dugan, T. R.; Cowley, R. E.; Brennessel, W. W.; Flaschenriem, C. J.; Bill, E.; Cundari, T. R.; Holland, P. L. J. Am. Chem. Soc. 2008, 130, 6624–6638.

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maintained at or below 1 ppm of O2 and H2O standard or by Schlenk techniques. Glassware was dried at 150 C overnight. NMR data were recorded on a Bruker Avance 500 or Avance 400 spectrometer at 22 C. All resonances in the NMR spectra are referenced to residual C6D5H at δ 7.16 ppm. In some cases, it was not possible to determine integrations because of overlap of signals. UV-vis spectra were measured on a Cary 50 spectrophotometer, using screw-cap cuvettes. Solution magnetic susceptibilities were determined by the Evans method.24 Elemental analyses were determined by Columbia Analytical Services, Tucson, AZ. Gas chromatography for detection of H2 used techniques reported recently.20,25 Pentane, diethyl ether, and toluene were purified by passage through activated alumina and “deoxygenizer” columns from Glass Contour Co. (Laguna Beach, CA). Deuterated benzene was first dried over CaH2, then over Na/benzophenone, and then vacuum transferred into a storage container. Before use, an aliquot of each solvent was tested with a drop of sodium benzophenone ketyl in THF solution. Celite was dried overnight at 200 C under vacuum. All reagents were purchased from Aldrich. CH3MgCl (3.0 M in diethyl ether) was used as received. (CH3)3SnF was prepared by a literature procedure.5 LtBuCoCH3 was prepared through a procedure that has been modified26 from the original procedure.6 Before use, LtBuCoCH3 was recrystallized from diethyl ether and washed with cold pentane to remove any possible impurities including LtBuH and trace Grignard reagent. Synthesis of LtBu2Co2F4Sn(CH3)2 (1). A resealable flask was charged with a brown solution of LtBuCoCH3 (945 mg, 1.64 mmol) in toluene (40 mL). (CH3)3SnF (660 mg, 3.60 mmol) was added, and the mixture was heated at 80 C for 24 h. The mixture was filtered and concentrated to 15 mL. Addition of diethyl ether (6 mL) and cooling at -40 C gave red crystals of LtBu2Co2F4Sn(CH3)2 3 2Et2O (796 mg, 65%). 1H NMR (C6D6): δ 63.8 (4H, p-aryl), 22.2 (8H, iPr methine or m-aryl), 20.8 (36H, tBu), -3.2 (24H, iPr methyl), -20.6 (2H, R-H), -27.2 (24H, iPr methyl), -38.0 (6H, Sn(CH3)2), -52.3 (8H, iPr methine or m-aryl) ppm. μeff (Evans, C6D6): 7.2(2) μB. Anal. Calcd for C80H132N4O2Co2F4Sn: C 64.29, H 8.90, N 3.75. Found: C 64.41, H 8.57, N 3.69. UV-vis (Et2O): 920 (4.4), 519 (11.3). IR (KBr): 2952 (s), 1459 (s), 1419 (s), 1372 (s), 1369 (s), 1328 (s), 1212 (m), 1156 (m), 1095 (m), 1059 (w), 792 (m), 756 (s) cm-1. Synthesis of LtBuCoF(py) (2). To a red solution of 1 (90 mg, 63 μmol) in diethyl ether (15 mL) was added pyridine (15 μL, 180 μmol) via microsyringe, leading to immediate formation of a white precipitate. After 1 h, the mixture was filtered, reduced to 8 mL, and cooled at -35 C to give light yellow crystals (54 mg, 72%). 1H NMR (C6D6): δ 123, 41, 32, 24, 19, 8.8, 4.7, 2.2, 1.2, 0.2, -1, -21, -37, -94 ppm. Assignments were not possible because of signal overlap and similar integrations of the peaks. μeff (Evans, C6D6): 4.7(4) μB. Anal. Calcd for C40H58N3CoF: C 72.92, H 8.87, N 6.38. Found: C 72.88, H 8.84, N 6.29. UV-vis (Et2O): 721 (0.11), 462 (0.51), 409(0.64). IR (KBr): 2959 (s), 1460 (m), 1423 (s), 1382 (s), 1363 (s), 1328 (s), 1213 (m), 1153 (m), 1095 (m), 1056 (w), 800 (m), 756 (s) cm-1. Reaction of LtBuCoF(py) (2) with Et3SiH to Form 3. In an argon-filled glovebox, a resealable NMR tube was loaded with 2 (7.3 mg, 9.7 μmol), Et3SiH (1.6 μL, 9.7 μmol), and C6D6 (0.3 mL). An internal reference consisting of a capillary containing a solution of LtBuFeCl (22.5 mM) in C6D6 was added. The tube was sealed and heated to 50 C for 7 h. Production of 3 was observed by 1H NMR spectroscopy with 86% yield. (24) (a) Baker, M. V.; Field, L. D.; Hambley, T. W. Inorg. Chem. 1988, 27, 2872–2876. (b) Schubert, E. M. J. Chem. Educ. 1992, 69, 62–62. (25) Yu, Y.; Brennessel, W. W.; Holland, P. L. Organometallics 2007, 26, 3217–3226. (26) Ding, K.; Pierpont, A. W.; Brennessel, W. W.; Lukat-Rodgers, G.; Rodgers, K. R.; Cundari, T. R.; Bill, E.; Holland, P. L. J. Am. Chem. Soc. 2009, 131, 9471–9472.

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Table 2. Details of X-ray Crystal Structures tBu

[L empirical formula fw cryst syst space group a (A˚) b (A˚) c (A˚) R (deg) β (deg) γ (deg) V (A˚3) Z F (g/cm3) μ (mm-1) R1, wR2 (I > 2σ(I)) GOF

Co(μ-F)2]2Sn(CH3)2 (1)

C80H132Co2F4N4O2Sn 1494.45 triclinic P1 12.426(1) 16.003(2) 21.112(2) 100.573(2) 91.507(2) 102.187(2) 4024.7(8) 2 1.233 0.769 0.0333, 0.0719 1.025

Independent Synthesis of LtBuCo(py) (3). In an argon-filled glovebox, a solution of KHBEt3 (49 mg, 356 μmol) in toluene (5 mL) was added to a solution of LtBuCoCl (211 mg, 354 μmol) in toluene (25 mL), causing the immediate formation of a yellow-brown solution of LtBuCo(μ-H)2CoLtBu, as reported separately.20 Pyridine (30 μL, 370 μmol) was added to the stirred mixture via microsyringe, giving a green mixture. After stirring for 30 min, the volatile materials were pumped off. The green residue was extracted with diethyl ether (15 mL) and filtered through Celite. The solution was reduced to 5 mL and cooled at -35 C to give green crystals (139 mg, 59%). 1H NMR (C6D6): δ 76.3 (2H, p-aryl), 44.2 (5H, pyridine), 33.9 (4H, m-aryl or iPr methine), 4.9 (12H, iPr methyl), 1.1 (12H, iPr methyl), -3.5 (18H, tBu), -52.8 (4H, m-aryl or iPr methine), -82.4 (1H, backbone) ppm. μeff (Evans, C6D6): 3.5(3) μB. Anal. Calcd for C40H58N3Co: C 75.09, H 9.14, N 6.57. Found: C 73.24, H 9.64; N 6.82. Repeated attempts at microanalysis did not yield better results, and it is possible that it cocrystallizes with NMR-silent impurities. UV-vis (Et2O): 432 (2.78), 403 (0.78). IR (KBr): 2959 (s), 2907 (m), 1542 (w), 1423 (m), 1461 (m), 1398 (m), 1362 (s), 1328 (m), 1210 (w), 1153 (m), 1095 (m), 1012 (w), 795 (m), 756 (s) cm-1. Electron Paramagnetic Resonance. EPR spectra were recorded with a Bruker ER200D spectrometer interfaced to an IBM PC for data recording. Data acquisition used a locally written program by Prof. Robert Kreilick at the University of Rochester. Spectra were recorded from 50 to 7150 G. A field modulation of 15 G and 100 kHz was used. The X-band microwave frequency was 9.415 GHz. The time constant was 100 ms. The sample was cooled to 10 ( 3 K using a liquid helium cryostat, and the temperature was calibrated by a thermocouple placed in a sample tube in the cavity. Magnetic Susceptibility Measurements. Magnetic susceptibility data were measured from powder samples of solid material in the temperature range 2-300 K using a SQUID susceptometer (MPMS-7, Quantum Design) with a field of 1.0 T. The experimental data were corrected for underlying diamagnetism by use of tabulated Pascal’s constants (χdia = -747  10-6 emu). (27) Available from http://ewww.mpi-muelheim.mpg.de/bac/logins/ bill/julX_en.php.

LtBuCoF(py) (2)

LtBuCo(py) (3)

C40H58CoFN3 658.82 triclinic P1 12.107(2) 18.244(3) 18.468(3) 89.740(2) 77.893(2) 75.628(2) 3859.0(10) 4 1.134 0.479 0.0630, 0.0866 0.989

C40H59CoN3 639.82 monoclinic P21/n 9.4374(15) 17.387(3) 21.780(3) 90 94.162(3) 90 3564.4(10) 4 1.194 0.512 0.0548, 0.1144 1.008

The susceptibility data, χT(T) or μeff(T), were simulated with our package julX for exchange-coupled systems written by E. B.27 The simulations are based on the spin-Hamilton operator suitable for the trimeric compound (1) with zero-field splitting for the cobalt ions:

^ ^ ^ ^ F ^ F ^ F F F F ^ ¼ -2J12 ½S 1 S 2  -2J13 ½S 1 S 3  -2J23 ½S 2 S 3  H 3 3 3 ^ F ^ ^ F F F F þ g1 μB S 1 3 B þ g1, 2 μB ðS 2 þ S 3 Þ 3 B X 2 þ DCo ½S^i, z -1=3SðS þ 1Þ i ¼2, 3

ð1Þ

where S1 = 1/2, S2 = 3/2, and S3 = 2 are spins for Sn(III), Co(II), and Co(III), respectively, g1 and g2,3 are the corresponding average electronic g value, and DCo is the axial zero-field splitting parameter for the cobalt ions. The magnetic moments in the simulations were obtained from the eigenfunctions |ψiæ of eq 1 by using the Hellman-Feyman theorem dEi/dB = Æψi|dH|dB|ψiæ. Powder summations were done by using a 16-point Lebedev grid. Attempted HDF of Fluoroarenes. In a representative experiment, a J. Young NMR tube was loaded with a solution of C6F6 (0.12 M), triethylsilane (0.12 M), precatalyst 2 (0.024 M), and solvent (C6D6 or THF-d8). The tube was heated to 45 or 80 C in an oil bath, and NMR spectra (19F, 1H) were recorded periodically. X-ray Crystallography. Key parameters are in Table 2, and details are in the Supporting Information.

Acknowledgment. This research was supported by the National Science Foundation (CHE-0112658), the A. P. Sloan Foundation, and the Department of Energy (DESC0001834). Supporting Information Available: Crystallographic data (CIF format). This material is available free of charge via the Internet at http://pubs.acs.org.