Titanium(IV) Trifluoromethyl Complexes: New Perspectives on

Article pubs.acs.org/Organometallics

Titanium(IV) Trifluoromethyl Complexes: New Perspectives on Bonding from Organometallic Fluorocarbon Chemistry Felicia L. Taw,† Aurora E. Clark,† Alexander H. Mueller,† Michael T. Janicke,† Thibault Cantat,† Brian L. Scott,† P. Jeffrey Hay,† Russell P. Hughes,*,‡ and Jaqueline L. Kiplinger*,† †

Los Alamos National Laboratory, Los Alamos, New Mexico 87545, United States Dartmouth College, 6128 Burke Laboratories, Hanover, New Hampshire 03755, United States

‡

S Supporting Information *

ABSTRACT: Trifluoromethyltrimethylsilane, (CH3)3SiCF3, in the presence of CsF serves as an excellent CF3 group-transfer reagent, and reaction with Cp2TiF2 in THF gives the titanocene trifluoromethyl fluoride complex Cp2Ti(CF3)(F) (1; Cp = C5H5) in 60% isolated yield. Reaction of complex 1 with the trimethylsilyl reagents, (CH3)3SiX (X = OTf = OSO2CF3, Cl, Br, I, N3, and OSO2Ph), in a tetrahydrofuran or toluene solution affords the corresponding Ti−CF3 derivatives Cp2Ti(CF3)(X) (X = OTf (2), Cl (12), Br (13), I (14), N3 (15), and OSO2Ph (16)) in good isolated yields of 67−84%. These compounds have been characterized by a combination of reactivity studies, IR and 1 H/13C{1H}/19F NMR spectroscopies, and single-crystal X-ray diffraction. The Ti−CF3 linkage in these complexes is remarkably robust, and although the α-C−F bonds are elongated, there is no evidence of an α-fluoride (Ti···F−CF2) between the electrophilic Ti(IV) metal center and any of the C−F bonds in the trifluoromethyl group in the solid-state or in solution. In the solid-state, these complexes are shock-sensitive; energetic decomposition of Cp2Ti(CF3)(F) (1) produces uniform spherical nanoparticles ranging from ∼70 to 120 nm in size and porous fluorinated oligomers and polymers containing both −(CF2−CF2)− and −(CF2−CFH)− units, as determined by a combination of SEM, XRD, XRF, XPS, and 19F MAS NMR. Density functional theory results show good agreement with experimental structural data obtained for Cp2Ti(CF3)(X) (X = F (1), OTf (2), Cl (12), N3 (15)) and accurately predicts longer Ti−CF3 distances than for each specific CH3 analogue, and the trend extends to structurally related Zr and Hf analogues. Simpler model compounds from groups 4 and 8 (M(CH3)4, M(CH3)3(CF3), M(CH3)3(CCl3), and M(CH3)3(CF2CF2CF2CF3); M = Ti, Zr, Hf, Fe, Ru, Os)) were also examined and show that, for group 4 complexes, π-bonding is a significant factor in shortening the strongly ionic M−CH3 relative to M−CF3, whereas for the predominantly covalent group 8 analogues, π-back-bonding helps to shorten the predominantly covalent M−CF3 relative to M−CH3. The bonding analysis suggests that the significant elongation of C−F bonds α to metals is mainly a consequence of the electropositivity of the group 4 metal centers, with minor, if any, contributions from π-effects; the bond-lengthening effect is most pronounced at the α-position and decays rapidly on moving away from the metal.

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decarbonylation of the parent trifluoroacetyl complex.4 Other trifluoromethyl complexes have been prepared in an analogous manner, but this technique requires the presence of a trifluoroacetyl ligand on a metal center that possesses a vacant coordination site to which the CF3 group can migrate. Moreover, thermolysis or photolysis is often necessary to promote decarbonylation.5 A second technique was introduced that involved oxidative addition of trifluoromethyl iodide to a lowvalent transition metal; however, this method was typically successful only for d8 or d10 metal substrates.5,6 A third route to perfluoroalkyl complexes consisted of metal atom reactions and radical chemistry, in which either the metal or the CF3 fragment was activated by an external high-energy source.5,7 More recently, the group 12 cadmium and mercury systems, Cd(CF 3) 2·2L

INTRODUCTION Continued interest in fluorocarbon chemistry stems from the unique properties that fluorinated materials exhibit.1 The high electronegativity of fluorine and great strength of C−F bonds conspire to give compounds that have low chemical and thermal reactivity, as compared with hydrocarbon analogues.2 As such, it is inevitable that these same characteristics of fluorocarbons present challenges to the preparation of fluorinated complexes. Specifically, although a vast number of transition-metal catalysts have been proven to affect the polymerization of hydrocarbon olefins,3 similar catalytic platforms do not exist for the polymerization of fluorinated olefins. As such, the discovery of new transition-metal perfluoroalkyl complexes and developing an understanding of their unusual behavior are crucial steps in this direction. In contrast to the arsenal of choices that are available for alkylation, only a handful of pathways exist to functionalize transition metals with perfluoroalkyl groups. The first transition-metal CF3 complex, Mn(CF3)(CO)5, was synthesized by thermal © 2012 American Chemical Society

Special Issue: Fluorine in Organometallic Chemistry Received: October 31, 2011 Published: January 27, 2012 1484

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complex (11) has also been formed from reaction of the corresponding Ce−H and (CH 3 ) 3 SiCF 3 . 21 Subsequent α-fluoride abstraction of a fluoride by the cerium metal center gives the cerium fluoride complex (1,2,4-tBu3-C5H2)2Ce-F and difluorocarbene. The Ti−CF3 complexes 1 and 2 were characterized by X-ray diffraction studies, which revealed that the Ti−C bond distances (2.221(3) Å for 1; 2.222(5) Å for 2)11 in the Ti− CF3 moiety were substantially longer than Ti−C bond distances (1.988−2.181 Å)22 reported for structurally related Ti−CH3 complexes. This directly contradicted existing tenets regarding M−C (M = middle- or late-row transition metal) bond length trends in M−CF3 complexes versus analogous M−CH3 complexes, in which a shorter M−C bond length is expected for the M−CF3 compounds.5 Indeed, structural comparisons of previously reported middle- and late-row transitionmetal perfluoroalkyl and corresponding alkyl compounds revealed a significant contraction (by ∼0.05 Å)5f of the metal−perfluoroalkyl bond length relative to the metal−alkyl bond length.5 For example, the M−CF3 bond distance in (C5Me5)Mo(CF3)(CO)323 was reported to be 2.248(5) Å, whereas in the related molybdenum complex, [η5-C5H4C( O)CH3]Mo(CH3)(CO)3,24 the M−CH3 bond distance was reported to be 2.303(5) Å. In the titanium case, an unexpected elongation of the metal−perfluoroalkyl (M−C) bond length was observed. Herein, we describe the synthesis, structural characterization, properties, and reactivity of a series of Ti−CF3 complexes and present density functional theory (DFT) calculations to investigate the bonding and electronic structure in these Ti−CF3 systems that gives rise to the unprecedented bond length trend in comparison to mid- and late-row transitionmetal systems.

(L = pyridine or 1,2-dimethoxyethane) and Hg(CF3)2, have been employed as CF3-transfer reagents for a variety of transitionmetal and main-group compounds.5,8 These reagents demonstrated high versatility and allowed for the preparation of many previously unknown transition-metal trifluoromethyl complexes. The techniques described above have allowed entry into the synthesis of M−CF3 complexes for virtually all the middle- and late-row transition metals. However, these same methods have failed for the early transition metals (groups 3, 4, and 5). An early report described the treatment of Ti(CH3)4 with trifluoromethyl iodide, which led to formation of Ti(CH3)3I.9 No Ti−CF3 species or any fluorine-containing compound could be detected. Researchers have also reported attempts to synthesize Zr−CF3 complexes using the aforementioned cadmium and mercury CF3 reagents, but only zirconium fluorides were isolated as the end products.10 Several years ago, we demonstrated that stable early transitionmetal perfluoroalkyl could be prepared with the synthesis and characterization of the titanium perfluoroalkyl complexes, Cp2Ti(CF3)(F) (1; Cp = C5H5) and Cp2Ti(CF3)(OTf) (2; OTf = OSO2CF3).11 Titanocene trifluoromethyl fluoride (1) was accessed by the addition of Ruppert’s reagent, (CH3)3SiCF3, to a tetrahydrofuran (THF) suspension of Cp2TiF2,12 and the titanocene trifluoromethyl triflate complex (2) was synthesized by treating 1 with the trimethylsilyl compound, (CH3)3SiOTf (eq 1).11

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RESULTS AND DISCUSSION Synthesis, Structural Characterization, and Reactivity. As shown in eq 2, the synthesis of titanium trifluoromethyl com-

Although widely used to trifluoromethylate organic compounds,13 Ruppert’s reagent has also been used in a few cases to transfer CF3 to late transition metals (Chart 1).14 For example, Chart 1. Examples of Metal Trifluoromethyl Complexes Prepared Using the CF3 Delivery Agent (CH3)3SiCF3

plexes was achieved by the addition of trimethylsilyl reagents, (CH3)3Si-X (X = OTf, Cl, Br, I, N3, and OSO2Ph), to a THF or toluene solution of Cp2Ti(CF3)(F) (1) to form Cp2Ti(CF3)(X) (X = OTf (2), Cl (12), Br (13), I (14), N3 (15), and OSO2Ph (16)). In all cases, the reactions were performed at room temperature, and release of (CH3)3SiF was observed by 1 H and 19F NMR spectroscopy as the reaction progressed. Complex 12 may be prepared in high yield (85−90% isolated) by addition of 5 equiv of (CH3)3SiCl to a THF solution of 1. After stirring at ambient temperature for 3 h, removal of volatile materials under reduced pressure gave an orange powder. Trace amounts of 1 were present in this crude product, which can be purified by crystallization from a toluene/hexanes solution at −30 °C. Alternatively, 5 equiv of (CH3)3SiCl may be added to

Caulton and co-workers showed that addition of (CH3)3SiCF3 to a ruthenium fluoride complex led to replacement of the fluoride ligand by CF3, giving 3.14a,b Subsequent α-migration of the fluoride from CF3 to the metal center resulted in formation of the final product, a ruthenium difluorocarbene fluoride complex (4). Since our initial communication, additional reports have shown that Ruppert’s reagent is quite effective as a CF3 delivery agent and gives a variety of stable first-row (Ni, 5;15 Cu, 616), second-row (Rh, 7;17 Pd, 818), and third-row (Pt, 9;19 Au, 1020) transition-metal complexes. A reactive cerium trifluoromethyl 1485

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Table 1. Selected 1H, 19F, and 13C{1H} NMR and IR Spectroscopic Data for Complexes 1, 2, and 12−16a Cp2Ti(CF3)X X X X X X X X a

= = = = = = =

F (1) N3 (15) Cl (12) Br (13) OSO2Ph (16) I (14) OSO2CF3 (2)

1

H [δ C5H5] 6.40 6.43 6.52 6.58 6.63 6.72 6.76

19

13

F [δ Ti−CF3] −24.0 −23.2 −21.2 −20.6 −24.5 −20.5 −25.6

153.1 153.9 153.2 155.2 152.4 159.3 151.8

C{1H} [δ CF3]

(q, (q, (q, (q, (q, (q, (q,

1

JC−F 1 JC−F 1 JC−F 1 JC−F 1 JC−F 1 JC−F 1 JC−F

= = = = = = =

382.6 381.8 382.1 381.5 383.1 380.0 382.9

13

C{1H} [δ C5H5]

Hz) Hz) Hz) Hz) Hz) Hz) Hz)

117.4 116.8 118.6 118.6 119.3 118.3 120.7

νTi−CF3

νC−Fsymm

692 689 691 690

1081 1065 1077 1074

689 694

1061 1082

Chemical shifts reported in parts per million; THF-d8 used as solvent. IR spectroscopic data are reported in cm−1.

have been prepared by reaction of the corresponding metal− trifluoromethyl complex with protic acids.5d However, addition of protic acids (HCl, HOTf, or HBAr'4 (Ar' = 3,5-(CF3)2-C6H3) or the methide-abstracting agent B(C6F5)3 to the Ti−CF3 complexes 1 and 2 led to immediate decomposition with the formation of intractable products. Selected 1H, 19F, and 13C{1H} NMR spectroscopic data for all Ti−CF3 complexes (1, 2, 12−16) are listed in Table 1. The 1 H NMR resonances for the Cp ligands fall in the range of δ 6.40−6.76 ppm. Within the series where X = halide, the Cp resonances shift downfield when going from X = F→Cl→Br→I; this is consistent with decreasing π donation from the halide to the d0 Ti metal center [i.e., X(π)→Ti(dπ) donation], with the fluoride ligand providing the greatest amount of donation to the metal center. This effect has been measured in numerous systems using NMR and IR spectroscopy, and the π-donor power of the halides has been ranked in the order F > Cl > Br > I, which is the opposite of what would be predicted based upon electronegativity arguments.26 The singlets that appear in the region of δ −20 to −25 ppm in the 19F NMR spectra and the quartets with large 1JC−F coupling constants (∼380−384 Hz) that appear in the region of δ 150−160 ppm in the 13C{1H} NMR spectra are diagnostic for the presence of CF3 groups. This is also consistent with the 19F and 13C{1H} NMR resonances for CF3 groups in other characterized middle- and late-row transition-metal trifluoromethyl complexes.5 Solid-state IR studies of complexes 1, 2, and 12−16 showed Ti−CCF3 stretching frequencies spanning 689−694 cm−1 and C−F (of Ti−CF3) stretching frequencies in the range of 1061− 1082 cm−1 (Table 1). The observed C−F stretches for the titanium compounds are characteristic of CF3 ligands bound to middle- and late-row transition metals (e.g., (CO)5MnCF3, νC−F = 1045, 1063 cm−1;27 Ru(CF3)Cl(CO)2(PPh3)2, νC−F = 1073, 1006 cm−1;28 Ir(CF3)(CO)2(PPh3)2, νC−F = 1088, 1005 cm−1;29 CpMo(CF3)(CO)3, νC−F = 1073, 1006 cm−1;30 CpFe(CF3)(CO)2, νC−F = 1068, 1042, 1015 cm−1).30 Typically, these frequencies are approximately 100 cm−1 lower than the mean C−F stretching frequencies in CF3X (X = Cl, Br, I; νC−F ∼ 1140 cm−1),31 indicating weaker C−F bonds in the metal complexes.27 X-ray diffraction studies confirmed the identity of the titanium trifluoromethyl complexes. Representative structures are provided for complexes 12 and 15 in Figure 1, and selected geometric parameters for complexes 1, 2, 12, and 15 are presented in Table 2. The titanium trifluoromethyl complexes all exhibit a bent-metallocene framework with a pseudotetrahedral coordination environment about the titanium metal center and the nonCp ligands lying within the metallocene wedge. For all structures, no α-fluoride interaction (Ti···F−CF2) is observed between the electrophilic titanium(IV) metal center and any of the C−F bonds in the trifluoromethyl group. Variable-temperature NMR spectroscopic studies (−80→+40 °C) on the complexes

a toluene solution of 1, allowed to stir for 3 h, and layered directly with hexanes to produce orange crystals of 12 in approximately 70% yield. Other Ti−CF3 derivatives of 1 may be prepared in an analogous manner. For complexes 2, 13, 14, and 16, use of 1 equiv of the appropriate trimethylsilyl reagent is sufficient to drive the ligand substitution reaction to completion. The bromidesubstituted complex 13 was formed (>95% by NMR spectroscopy) within 45 min at room temperature, whereas the triflate (2), iodide (14), and benzenesulfonate (16) analogues were formed within seconds under the same conditions. Synthesis of the azide complex (13) required the addition of 10 equiv of (CH3)3SiN3 to a THF solution of 1 and a reaction time of 2 h. Isolated yields for compounds 2 and 13−16 varied between 67 and 84%. Addition of other trimethylsilyl reagents (CH3)3Si-R (R = CH3, C6H5, CO2Me, CCH, Si(CH3)3, CH2CN, NCO, and C6F5) or Ph3SiH to 1 resulted in no reaction (by 1H and 19F NMR spectroscopy), even upon heating the reaction mixture to 60 °C for several days. Further, numerous attempts to prepare the bis(trifluoromethyl) complex, Cp2Ti(CF3)2, by treating Cp2TiF2 or 1 with (CH3)3SiCF3 or Zn(CF3)2·2(pyridine) were also unsuccessful. This contrasts behavior of late metal iron(II) and nickel(II) fluoride complexes, which display reaction chemistry with these trimethylsilyl reagents to give complexes, such as (LtBu)Fe-CCSiMe3 (LtBu = N,N′-bis(2,6-diisopropylphenyl)2,2-6,6-tetramethylheptane-3,5-diiminate)25 and most likely (dippe)Ni(CF3)2 (dippe = 1,2-bis(diisopropylphosphino)ethane).15 This marked difference in reactivity is likely a reflection of the stronger Ti−F bond compared with the later transition metal− fluoride bonds. The Ti−CF3 complexes 1, 2, and 12−16 decompose to intractable products at room temperature over the course of several hours (1, 2, 12−16) to several days (12) and require storage under cold conditions (−30 °C). In particular, complexes 1, 2, and 12 are shock-sensitive in the solid state (vide infra) and should be handled with great care. In solution, the stability of all the Ti−CF3 compounds is enhanced by the addition of Lewis bases, such as THF, pyridine, 4-(N,N-dimethylamino)pyridine, P(OMe)3, or phosphines (PPh3, PMe3). Solutions of complexes 2, 14, and 16 require the presence of a Lewis base, such as pyridine, to stabilize these species at room temperature and allow characterization and reactivity studies. Because complex 1 is stable, the reaction between Cp2Ti(CO)2 and CF4 was also examined, but no reaction was observed at room temperature or with heating at 50 °C even after 5 days. No reaction was observed between any of the Ti−CF3 complexes 1, 2, and 12−16 and nucleophilic reagents, such as nitriles, isocyanides, amines, carbon monoxide, or benzophenone. Addition of olefins, fluorinated olefins, diphenyldiazomethane, and diphenylacetylene also resulted in no reaction. A variety of middle- and late-row transition-metal dihalocarbenes 1486

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Figure 1. Molecular structures of complexes 12 (left) and 15 (right) with thermal ellipsoids projected at the 25% probability level. Hydrogen atoms have been omitted for clarity.

Table 2. Selected Metrical Parameters for Complexes 1, 2, 12, and 15 1 (X = F) Ti−X (Å) Ti−CF3 (Å) C−F (Å)

C−Ti−X (deg) F−C−F (deg)

Ti−C−F (deg)

1.831(2) 2.221(3) 1.362(3) 1.365(4) 1.370(3) 89.35(9) 103.7(2) 102.3(2) 102.0(2) 114.0(2) 114.5(2) 118.4(2)

2 (X = OTf) 1.985(4) 2.222(5) 1.367(6) 1.362(4) 1.362(4) 88.3(2) 103.3(3) 103.3(3) 102.8(4) 114.2(2) 114.2(2) 117.2(3)

1.317(4) 1.314(4) 1.314(4)

corroborate the solid-state data and confirm the absence of α-fluoride interactions in solution as no change in the CF3 19F NMR shift or 1JC−F was observed. Comparisons for the metrical parameters of the Ti−CF3 moiety can only be made to known middle and late transitionmetal trifluoromethyl complexes. As in the case of complexes 1 (2.221(3) Å) and 2 (2.222(5) Å), the metal−carbon bond distances in the Ti−CF3 moiety for 12 (2.301(6) Å) and 15 (2.239(1) Å) are longer than Ti−C bond distances (1.988− 2.181 Å) reported for structurally related Ti−CH3 complexes.22 For the trifluoromethyl group in complexes 1, 2, 12, and 15, the average C−F bond length (1.357(4) Å) and the average F− C−F bond angle (103.7(3)°) are quite similar to values found in known middle- and late-row transition-metal CF3 complexes that are thermally stable.5 Although direct quantitative comparisons are difficult, these observations are somewhat surprising as the Ti−CF3 complexes display low thermal stability. As complex 2 contains a triflate (OTf = OSO2CF3) group, it provides a nice internal platform to compare the metrical parameters of the C−F bonds on the M−CF3 group with those on the coordinated triflate ligand. Examination of salient bond lengths between the titanium complexes reveals statistically relevant lengthening of the C−F bond distances for the Ti−CF3 groups in 1, 2, 12, and 15 (C−Fave = 1.357(4) Å) compared with the 1.314(4)−1.317(4) Å for the CF3 moiety in the OTf ligand in complex 2. This suggests that the electropositive titanium(IV) metal center is siphoning away electron

(OTf) (OTf) (OTf)

12 (X = Cl)

15 (X = N3)

2.335(2) 2.301(6) 1.369(7) 1.306(6) 1.286(6) 89.27(13) 105.1(5) 104.7(5) 108.9(5) 110.2(4) 113.3(4) 114.0(4)

1.991(1) 2.239(1) 1.379(2) 1.382(2) 1.376(2) 89.59(5) 103.4(1) 102.5(1) 102.9(1) 113.3(1) 115.2(1) 117.7(1)

density from the Ti−CCF3 linkage, resulting in a corresponding lengthening of the trifluoromethyl C−F bonds. The observed lengthening of both the Ti−CCF3 and the C−F bonds coupled with the IR data points to a weakening of these C−F bonds. The fact that the d0 titanium metal center does not disassemble the CF3 group is remarkable. The Ti−F (1) and Ti−OOTf (2) bond distances of 1.831(2) and 1.985(4) Å, respectively, are somewhat shorter than the average Ti−F (1.848 Å) and Ti−OOTf (2.036 Å) bond distances reported for similar titanium(IV) terminal fluoride32 and monotriflate33 complexes. Likewise, for the new structure 12 presented in this work, the Ti−Cl bond length (2.335(2) Å) in complex 12 is slightly contracted when compared to other Ti−Cl bond distances (2.346−2.379 Å) in similar titanocene chloride complexes.34 A comparison of the Ti−N distances in the titanium azide complex 15 versus the known bis(azide) complex, Cp2Ti(N3)2 (17),35 also shows that the Ti−N distance (1.991(1) Å) in 15 is smaller than the Ti−N distance (2.03(1) Å) reported for 17. This consistent bond shortening across the series of Ti−X complexes (where X is the α-heteroatom F (1), O (2), Cl (12), or N (15)) is likely due to the electron-withdrawing effect of the CF 3 ligand, which increases the electrophilicity of the Ti center. However, put another way, this bond shortening also suggests Ti−X multiple bond character in these complexes resulting from intramolecular X→Ti π-donation, with consequent transfer 1487

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Figure 2. SEM images of the two morphologies of the fluorinated material formed upon the energetic decomposition of Cp2Ti(CF3)(F) (1): (a) spherical nanoparticles and (b) highly porous foam.

−(CF2−CF2)− and −(CF2−CFH)− units (see Figures 3 and 4). The XPS spectrum was scaled using the lowest binding energy peak of carbon to be adventitious, with the higher binding energy

of electron density to the d0 titanium metal center to stabilize the Ti−CF 3 interaction. 26b Finally, the N−N distances in the azide ligand of 15 are 1.219(2) and 1.142(2) Å for N(1)−N(2) and N(2)−N(3), respectively. Corresponding N−N distances in the bis(azide) complex (17) are slightly contracted (1.18(2) and 1.10(2) Å, respectively) in comparison. Single-crystal X-ray diffraction studies were also performed on the bromide (13) and iodide (14) complexes, but these structures were plagued with significant disorder between the CF3 and halide ligands and only allowed assignment of connectivity.36 Solid-State Stability of Titanocene Trifluoromethyl Complexes. In the solid-state, complexes 1, 2, and 12 are shock-sensitive and should be handled with great care in amounts no larger than 0.150 g. Complex 1 is even shocksensitive in hexanes. The instability of the Ti−CF3 systems is not entirely unexpected given the facile α- and β-fluoride elimination pathways available to early metal fluoroalkyl complexes to give high lattice energy metal fluorides. In fact, such favorable decomposition pathways have been used to explain previous failures to prepare perfluoroalkyl complexes similar to 1, 2, and 12−16. Energetic decomposition of Cp2Ti(CF3)(F) (1) was accomplished by tapping the solid with a glass pipet tip (inside a 20 mL scintillation vial). The resulting decomposition produces intractable solids. The volatiles were analyzed by 19F NMR, which determined that no CF4 was produced in the decomposition. The decomposition products were collected on a glass slide using two different methods and viewed using scanning electron microscopy (SEM), which showed the formation of uniform spherical nanoparticles ranging from ∼70 to 120 nm in size (Figure 2a) and a highly porous foam (Figure 2b). The spherical nanoparticles were obtained by collecting the volatile products onto a slide mounted directly above the Cp2Ti(CF3)(F) (1) as it energetically decomposed. The porous foam was generated by allowing the decomposition of complex 1 to take place within an opened scintillation vial, followed by pouring the resulting vapor (decomposition products) onto a glass slide. X-ray fluorescence (XRF) measurements indicated that residual titanium was present in the decomposition products. The absence of a diffraction pattern using powder X-ray diffraction (XRD) revealed that the spheres are amorphous. The porous foam was determined by both X-ray photoelectron spectroscopy (XPS) and 19F magic-angle spinning (MAS) NMR to be a mixture of fluorinated oligomers and polymers containing both

Figure 3. Wide-scan survey spectrum (XPS) of the porous foam produced from the energetic decomposition of Cp2Ti(CF3)(F) (1) detecting that the material contains fluorine and carbon.

Figure 4. 19F magic-angle spinning (MAS) NMR spectra of the porous foam produced from the energetic decomposition of Cp2Ti(CF3)(F) (1) (top), polytetrafluoroethylene tape (middle), and a blank sample (bottom).

peaks being assigned to F-bonded carbon atoms. The F signal also showed binding energy shifts consistent with fluorocarbon 1488

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oligomers and polymers. Interestingly, a comparison of the 19F MAS NMR spectrum obtained for the porous foam with that of PTFE reveals that Teflon-like materials are also present in this mixture. To the best of our knowledge, the energetic decomposition of a fluorinated organometallic complex represents a new method for producing fluorinated oligomers and polymers. Computational Study. Comparison of M−CCF3 and M− CCH3 Bond Lengths. Several different bonding rationales have been proposed to account for the M−CCF3 bond length contraction observed in middle- and late-row transition-metal perfluoroalkyl complexes. One early suggestion invoked π-back-bonding from the metal dπ electrons to the low-lying antibonding σ* C−F orbitals, thus imparting some double-bond character to the M−CCF3 bond and shortening it relative to the analogous M−C CH3 bond in a metal alkyl complex. 37 A completely analogous description presented this phenomenon in terms of hyperconjugation within the M−CF3 moiety, leading to a shortened M−C bond due to partial contribution from doubly bonded metal−carbon [MCF2]+[F−] resonance structures.30 Later studies seemed to rule out such π-effects and accounted for the metal perfluoroalkyl bond length contraction using arguments strictly within the σ-bonding regime. A carbonyl force constant analysis on (CO)5MnR (R = CH3 and CF3) separated the σ- and π-bonding effects and concluded that the CH3 and CF3 ligands are equally poor π-electron acceptors.38 Calculations by Hall and Fenske on the same Mn compounds indicated that the highly electropositive nature of the fluorinated carbon, as well as the enhanced s character of the CF3 group within the M−CCF3 σ-orbital, concentrates the electron density closer to the central carbon and leads to the observed M−C bond length shortening.39 Intuitively, this is consistent with Bent’s Rule,40 which predicts a metal− fluoroalkyl bond length contraction based upon the high electronegativity of fluorine. This leads to bond formation with carbon centers that deviate from sp3 hybridization such that minimal s character is present in the C−F bond and a larger amount of s character contributes to metal−carbon (M−CCF3) bonding. The M−CCF3 bond is then contracted relative to its alkyl counterparts because the spatial extent of the 2s orbital is less than that of the 2p. These arguments appear reasonable assuming predominantly covalent bonding, but it seems clear that the s character in the carbon orbital should be less important in determining bonding distance in a more ionic interaction. For the last three decades, these explanations have adequately accounted for the metal perfluoroalkyl bond length contraction observed in middle- and late-row transition-metal species. Certainly, M−C bond length shortening has been a consistent trend in all of the previously reported middle- and late-row M−CF3 complexes that have been structurally characterized, from group 6 Cr41 and Mo23 to group 10 Ni15 and Pt19 complexes. However, the early transition-metal Ti−CF3 complexes (1, 2, 12, 15) described above follow the opposite trend, exhibiting metal perfluoroalkyl bond length elongation when compared with analogous Ti−CH3 complexes. Almost three decades ago, Oberhammer discussed the differences in structural parameters between CH3 and CF3 groups bound to a variety of main group elements E and pointed out that the charge distribution within the CX3 group correlated well with the bond length to E.42 At that time, simple charge distribution calculations showed that, in a CH3 group, the carbon bears a negative charge, with hydrogens partially positive, whereas in CF3, the carbon is strongly positive and the fluorines

negatively charged. When attached to electronegative (δ−) elements, the bond to CH3 was shorter than that to CF3, whereas for strongly δ+ elements, the reverse trend was observed. Such arguments might also be applied to analogous transition-metal compounds in which electronegativity increases from left to right in a particular row and could explain why bonds from CH3 and CF3 to the electropositive Ti described here show the opposite trend compared with analogues to the right in the periodic table. To probe this possibility, DFT calculations were carried out on the compounds reported here, Cp2Ti(CF3)(X) [X = F (1), Cl (12), Br (13), I (14), N3 (15), and OTf (2)]; their CH3 analogues, Cp2Ti(CH3)(X) [X = F (18), Cl (19), Br (20), I (21), N3 (22), and OTf (23)]; and their zirconium and hafnium halide analogues, Cp2Zr(CR3)(X) (R = H, X = F (26), Cl (27), Br (28), I (29); R = F, X = F (30), Cl (31), Br (32), I (33)) and Cp2Hf(CR3)(X) (R = H, X = F (36), Cl (37), Br (38), I (39); R = F, X = F (40), Cl (41), Br (42), I (43)), respectively, in order to compare and understand the differences in bonding as a function of both the metal electronic structure and the choice of CF3 or CH3 ligand. Some simpler model compounds from groups 4 and 8 have also been examined to gain greater insight (vide infra). In addition, properties of some selected titanium and zirconium compounds containing CCl3 ligands, Cp2M(CCl3)(X) (M = Ti, X = F (24), Cl (25); M = Zr, X = F (34), Cl (35)), were computed. Density functional theory (DFT) calculations were carried out using the B3LYP functional43 and the triple-ζ LACV3P**+ + basis set,44 as implemented in the Jaguar suite of programs. A comparison of the metrical parameters for the Cp2Ti(CF3)(X) compounds (1, 2, 12, and 15) is presented in Table 3. Good agreement is observed between the measurements obtained experimentally and those calculated by DFT, allowing enhanced credence to the DFT-calculated metrics for the additional hypothetical complexes discussed. No Cp2Ti(CH3)(X) compounds have been crystallographically characterized,45 so in order to further appraise the ability of our chosen method to predict Ti−CH3 metrics, we calculated the structures of Cp2M(CH3)2 (M = Ti (44), Zr (46), Hf (48)), the structures of which have all been crystallographically determined.22c,46 Results are summarized in Table 4 and once again show excellent agreement between DFT at this level and crystallographically determined values. Significantly, for the Cp2Ti(CF3)(X) compounds (Table 3), DFT predicts longer Ti−CCF3 distances than for each specific CH3 analogue, and the trend extends to the Zr and Hf analogues. However, it is noteworthy that, within 0.01 Å, the Ti−CCH3 distances are essentially constant at 2.18 Å, Ti−CCF3 at 2.25 Å, Zr−CCH3 at 2.30 Å, Zr−CCF3 at 2.38 Å, Hf−CCH3 at 2.26 Å, and Hf−CCF3 at 2.35 Å, whereas the M−X distances vary significantly. The crystallographic distance values for the Ti complexes are likewise essentially the same. Identical values (to within 0.01 Å) and trends between metals are observed for Cp2M(CH3)2 (M = Ti, Zr, Hf) (Table 4). There seems to be no significant effect of the ancillary ligand on the M−C bond lengths for CH3 or CF3. The quantum theory of atoms in molecules (QTAIM) has been used to determine characteristics of the electron density at the bond critical points (BCP) for these molecules.47 A number of characteristic parameters can be calculated at the BCP (r) and have been used to elucidate the nature of the bonds between atoms in molecules. The electron density [ρ(r)] and its Laplacian [∇2ρ(r)] have been used extensively to characterize 1489

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Table 3. Metric, NBO, NPA, and QTAIM Data for the Group 4 Complexes Cp2M(CR3)(X) (M = Ti; R = H, F; X = F, Cl, Br, I, N3, CF3SO3. M = Ti, Zr; R = Cl; X = F, Cl. M = Zr, Hf; R = H, F; X = F, Cl, Br, I) distance (Å) compound Cp2Ti(CH3)(F) (18) Cp2Ti(CH3)(Cl) (19) Cp2Ti(CH3)(Br) (20) Cp2Ti(CH3)(I) (21) Cp2Ti(CH3)(N3) (22) Cp2Ti(CH3) (OTf) (23) Cp2Ti(CF3)(F) (1) Cp2Ti(CF3)(Cl) (12) Cp2Ti(CF3)(Br) (13) Cp2Ti(CF3)(I) (14) Cp2Ti(CF3)(N3) (15) Cp2Ti(CF3)(OTf) (2) Cp2Ti(CCl3)(F) (24) Cp2Ti(CCl3)(Cl) (25) Cp2Zr(CH3)(F) (26) Cp2Zr(CH3)(Cl) (27) Cp2Zr(CH3)(Br) (28) Cp2Zr(CH3)(I) (29) Cp2Zr(CF3)(F) (30) Cp2Zr(CF3)(Cl) (31) Cp2Zr(CF3)(Br) (32) Cp2Zr(CF3)(I) (33) Cp2Zr(CCl3)(F) (34) Cp2Zr(CCl3)(Cl) (35) Cp2Hf(CH3)(F) (36) Cp2Hf(CH3)(Cl) (37) Cp2Hf(CH3)(Br) (38) Cp2Hf(CH3)(I) (39) Cp2Hf(CF3)(F) (40) Cp2Hf(CF3)(Cl) (41) Cp2Hf(CF3)(Br) (42) Cp2Hf(CF3)(I) (43) a b

M−CR3 2.176

M−X 1.848

angle (deg) X−M−CR3 93.4

atomic NPA charges qM

qX

qCR3

0.94

−0.52

−0.87

qCR3

group charge

bond ionicity

C→M hybrid

CR3

M−CR3

CR3

0.20

−0.27

28%

QTAIM BCP data (M−C) ρ(r)

H(r)

sp

2.9

0.094

−0.038

2.9

0.094

−0.038

2.182

2.350

92.1

0.57

−0.30

−0.85

0.20

−0.25

26%

sp

2.180

2.592

91.9

0.52

−0.27

−0.86

0.21

−0.23

32%

sp2.8

0.094

−0.038

2.180

2.821

91.7

0.50

−0.23

−0.87

0.21

−0.24

28%

sp2.8

0.094

−0.038

2.183

1.999

92.3

0.71

−0.45

−0.86

0.20

−0.26

28%

sp2.8

0.093

−0.037

2.172

1.986

91.5

0.84

−0.88

−0.85

0.21

−0.22

25%

sp2.9

0.095

−0.039

2.248 2.221(3)a 2.257 2.301(6)b 2.257

1.836 1.831(2)a 2.320 2.335(2)b 2.553

92.0 89.35(9)a 91.3 89.27(13)b 91.5

0.86

−0.50

0.80

−0.38

−0.34

36%

sp1.6

0.080

−0.026

0.47

−0.25

0.82

−0.38

−0.32

38%

sp1.6

0.079

−0.026

0.41

−0.20

0.81

−0.38

−0.33

38%

sp1.6

0.079

−0.026

2.263

2.777

91.6

0.38

−0.14

0.81

−0.38

−0.33

40%

sp1.5

0.078

−0.025

2.252 2.239(1)b 2.245 2.222(5)a 2.319

1.972 1.991(1)b 1.958 1.985(4)a 1.834

90.4 89.59(5)b 90.3 88.3(2)a 90.9

0.61

−0.44

0.81

−0.38

−0.33

36%

sp1.4

0.079

−0.025

0.76

−0.86

0.81

−0.38

−0.33

34%

sp1.6

0.081

−0.027

0.93

−0.49

−0.42

0.01

−0.39

50%

sp1.5

0.069

−0.022

2.333

2.322

91.4

0.53

−0.24

−0.41

0.01

−0.38

52%

sp1.5

0.067

−0.021

2.298

1.992

98.5

1.82

−0.64

−1.10

0.21

−0.47

52%

sp2.4

0.088

−0.024

2.295

2.466

97.0

1.42

−0.39

−1.10

0.21

−0.47

50%

sp2.4

0.089

−0.024

2.3

0.089

−0.024

2.295

2.674

97.2

1.32

−0.31

−1.10

0.21

−0.47

50%

sp

2.294

2.896

97.6

1.25

−0.25

−1.10

0.21

−0.47

52%

sp2.3

0.089

−0.024

2.384

1.982

98.6

1.74

−0.63

0.66

−0.38

−0.48

55%

sp1.4

0.073

−0.014

2.383

2.443

96.5

1.33

−0.36

0.67

−0.38

−0.47

54%

sp1.4

0.073

−0.014

2.383

2.646

96.9

1.22

−0.27

0.67

−0.38

−0.47

54%

sp1.4

0.073

−0.014

2.386

2.858

97.2

1.14

−0.18

0.67

−0.38

−0.47

54%

sp1.3

0.073

−0.014

2.427

1.980

98.0

1.81

−0.63

−0.57

0.01

−0.54

69%

sp1.2

0.067

−0.012

2.427

2.445

96.8

1.39

−0.35

−0.55

0.01

−0.52

68%

sp1.2

0.067

−0.012

2.268

1.953

97.2

2.01

−0.65

−1.15

0.21

−0.52

58%

sp2.2

0.094

−0.028

2.266

2.434

95.9

1.63

−0.40

−1.13

0.21

−0.50

56%

sp2.1

0.095

−0.029

2.265

2.645

95.7

1.52

−0.32

−1.13

0.21

−0.50

56%

sp2.0

0.095

−0.029

2.263

2.864

95.9

1.45

−0.25

−1.14

0.22

−0.48

57%

sp2.0

0.095

−0.029

2.355

1.944

96.7

1.93

−0.64

0.62

−0.38

−0.52

60%

sp1.4

0.078

−0.018

2.354

2.410

95.2

1.53

−0.37

0.63

−0.38

−0.51

58%

sp1.3

0.079

−0.018

2.354

2.613

95.4

1.42

−0.28

0.63

−0.38

−0.51

58%

sp1.3

0.079

−0.018

1.32

−0.19

0.63

−0.38

−0.51

58%

1.2

0.078

−0.018

2.356

2.826

95.8

sp

X-ray structures of Cp2Ti(CF3)(F) and Cp2Ti(CF3)(OTf): Taw, F. L.; Scott, B. L.; Kiplinger, J. L. J. Am. Chem. Soc. 2003, 125, 14712−14713. This work. 1490

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2.360 Cp2Hf(CH3)(CF3) (51)

X-ray structure of Cp2Ti(CH3)2: Thewalt, U.; Wöhrle, T. J. Organomet. Chem. 1994, 464, C17−C19. bX-ray structures of Cp2Zr(CH3)2 and Cp2Hf(CH3)2: Hunter, W. E.; Hrncir, D. C.; Bynum, R. V.; Penttila, R. A.; Atwood, J. L. Organometallics 1983, 2, 750−755.

0.62 0.22 −1.17 1.77

0.21 −1.16 1.85

96.3 95.8(5)b 96.4 2.269 2.237(12)b 2.255

the concentration or depletion of electron density at a BCP as being indicative of electron-sharing (covalency) or closed-shell interaction (ionic or nonbonded) at this point.48 It has been noted that Laplacians may not be the most reliable indicators for some polar bonds,49 but Cremer and Kraka have suggested that the sum [H(r)] of the kinetic energy density [G(r)], and the potential energy density [V(r)] at the BCP is a more reliable indicator.50 Dominance of the potential energy density at the BCP gives a negative value for H(r) and is indicative that electron density buildup at the BCP is stabilizing, suggestive of covalency; a positive value of H(r) indicates the dominance of the kinetic energy density, in which electron density accumulation at the BCP is destabilizing, typical of a closed-shell interaction. Typical values of H(r) for covalent bonds lie between −0.15 and −0.45 au, and clearly, the best direct comparisons are between like pairs of atoms.50a,b Thus, in Cp2Ti(CH3)(F), the value of H(r) at the C−H BCPs of the CH3 ligand averages −0.237, clearly indicative of covalency, whereas that at the Ti−F bond is −0.008, indicative of a much more ionic character in the latter; similarly, H(r) at the C−F BCP in Cp2Ti(CF3)(F) is −0.280, whereas that at the Ti−F BCP is −0.010. Values of ρ(r) and H(r) at the M−C bond critical points are presented in Tables 3 and 4. The natural bond orbital (NBO)51 method allows bond ionicities to be calculated from the NBO polarization coefficients (c) for the M−CX3 bonds according to the expression

ionicity MC =

c M2 − cC2 c M2 + cC2

and these appear in Tables 3 and 4. Once again, the ionicities are reasonably constant within a given series, with each M−CF3 being slightly more ionic than the corresponding M−CH3 analogue. The greatest difference in bond ionicity between M−CH3 and M−CF3 is observed for Ti, with a general increase in bond ionicity for both ligands on descending the group. The bond ionicity trends are consistent with those observed in the H(r) values. Charges (q) on the metal centers and the individual atoms of the CR3 ligands were calculated using the NBO/NPA method,51b,c which is suggested to be among the better methods for apportioning charge within molecules,51c and are also presented in Tables 3 and 4, along with group charges for the CR3 ligands. For the sake of comparison, the corresponding atom NPA charges and lone-pair hybridizations for the CR3− anions were also calculated at the same DFT level: CH3− qC = −1.40, qH = +0.13 (sp5.4 hybrid lone pair), CF3− qC = +0.47, qF = −0.49 (sp0.6 hybrid lone pair), and CCl3− qC = −0.34, qCl = −0.22 (sp0.3 hybrid lone pair). In agreement with Oberhammer’s results,42 the carbon of the CH3− anion bears a strong negative charge, whereas that of the CF3− anion bears a significant positive charge, with large negative charges on the F atoms. The lone pair of the CH3− anion resides in an orbital with large p character, whereas that of the CF3− anion has considerably greater s character, as predicted by Bent’s Rule. The CCl3− anion is interesting in that its charges are distributed in the same sense as the CH3− anion, with a negatively charged carbon, but, like the CF3− anion, the lone pair is in a high s-character orbital. The NPA charges in CF3− and CCl3− anions agree closely with those reported by Schrobilgen.52 Examination of Tables 3 and 4 illustrates that the trends in charge distribution within the CR3 ligands are maintained in all Ti, Zr, and Hf compounds. Remarkably, the charges on the

a

−0.029 sp2.0 −0.51 −0.38

−0.53

−0.52

60%

60%

60%

sp1.9

sp1.2

0.076

−0.017

0.096

−0.027 0.093

−0.025 0.090 −0.011 0.065 sp1.0 sp2.2 54% 0.00 0.22 −1.10 99.7 2.287 2.432

1.61

0.22

Cp2Ti(CH3)(CF3) (45) Cp2Ti(CH3)(CCl3) (46) Cp2Zr(CH3)2 (47)

Cp2Ti(CH3)2 (44)

Cp2Zr(CH3)(CF3) (48) Cp2Zr(CH3)(CCl3) (49) Cp2Hf(CH3)2 (50)

2.394

1.56

−1.11

0.65

−0.44

−0.56

70%

−0.024 56% 54% −0.49 −0.38

−0.45

54% −0.47 0.21 −1.10

99.0 95.6(12)b 100.8 2.300 2.277(5)b 2.287

1.63

0.21 89.8 2.159 2.321

0.76

−0.89

−0.56

sp

sp2.2

sp1.3

0.070

−0.013

0.089

−0.023 0.087

−0.040 0.098 −0.020 0.067

2.2

sp1.3 sp2.8 30% −0.46 −0.26 −0.01 −0.43

0.78 0.21 −0.91 2.257

0.70

0.21 −0.91

91.8 91.3(1)a 89.7 2.170 2.175(2)a 2.155

0.77

qCH3 compound

M− CR3

M−CH3

C−M−C

qM

qCH3

qCR3

54%

−0.040 −0.024 0.078 sp1.5 sp2.7 32% −0.28 −0.39

−0.28

−0.39

40%

sp2.5 34%

M− CR CR3 qCR3

CH3

Article

0.098

−0.038 0.095

ρ(r) M− CH H(r) M− CR M− CR M− CH M− CH

ρ(r) M− CR

QTAIM BCP data (M−C) C → M hybrid bond ionicities group charge atomic NPA charges angle (deg) distances (Å)

Table 4. Metric, NBO, and NPA Data for the Group 4 Complexes Cp2M(CH3)2, Cp2M(CH3)(CF3) (M = Ti, Zr, Hf) and Cp2M(CH3)(CCl3) (M = Ti, Zr)

H(r) M− CH

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giving a 12-valence electron maximum, after which additional coordination must be accommodated by multicenter bonding motifs.51c Therefore, for the complexes Cp2Ti(X)(Y) described in Tables 3 and 4, the necessity of a 16-valence electron count makes a detailed NBO analysis more complex. To reduce steric effects and to gain more insight into σ- and π-effects, all of which contribute to differences in bond lengths, a model system designed to accommodate a maximum of 12 valence electrons was studied. The tetrahedral d0 compounds M(CH3)4 (M = Ti (52), Zr (55), Hf (58)) are well known, and their structures have been studied computationally.55 The NBO description of these molecules involves a reference Lewis structure with 4sd3 hybrids accommodating the 2-center/2-electron σ-bonds to CH3 in a valence saturated structure, leaving two empty, degenerate d orbitals. Similarly, the hypothetical d4 compounds M(CH3)4 (M = Fe (61), Ru (64), Os (67)) represent a set of isostructural and isovalent analogues, with the same NBO Lewis structure for their singlet states, except with the two degenerate d orbitals filled. Thus, in the group 4 compounds, the metals are only capable of being π-acceptors, whereas in the group 8 analogues, they are only capable of being π-donors. Accordingly, the structures of M(CH3)4 (M = Ti (52), Zr (55), Hf (58), Fe (61), Ru (64), Os (67)), M(CH3)3(CF3) (M = Ti (53), Zr (56), Hf (59), Fe (62), Ru (65), Os (68)), and M(CH3)3(CCl3) (M = Ti (54), Zr (57), Hf (60), Fe (63), Ru (66), Os (69)) were optimized, as singlets, using the identical functional and basis set as before. The calculated Kohn−Sham LUMO for Zr(CH3)4 (55) and HOMO Ru(CH3)4 (64) are shown in Figure 5, along with the corresponding NBOs. All the compounds were also subjected to NBO, NPA, and QTAIM analyses. Results are presented in Table 5.

hydrogen atoms of every CH3 ligand remain essentially constant at +0.21 for all compounds, independent of the metal or the anionic ligand; similarly, the CF3 fluorine atoms bear an essentially constant charge of −0.38 throughout and the Cl atoms in the CCl3 examples are constant at ∼0.01. The idea of the different charge capacities of atoms in molecules has been discussed in detail by Politzer and Murray.53 Inspection of the charges on the carbon atoms is also instructive. For the CH3 and CCl3 ligands, the charge on carbon becomes more negative on going from Ti→Zr→Hf, but for each metal, it is independent of the anionic ligand. Likewise, the CF3 carbon becomes less positive on going from Ti→Zr→Hf, but again remains constant for a particular metal. In contrast, the negative charges on ancillary ligand X are significantly different and the positive charge on the metal center also varies significantly within each series. The carbon hybrid orbital in the M−CR3 bond depends strongly on R, in the same manner as in the CR3− anions, with significantly higher s character for R = F and Cl; there is a small dependence on M, with slight increases in s character in each class on descending the group. The hybridization is essentially independent of the ancillary ligand in each subgroup. Bond ionicities for M−CR3 depend on R, decreasing as Cl > F > H, and increase somewhat on descending the group, with only a small dependency on ancillary ligand X for each M. However, the electron densities ρ(r) and the total energy densities, illustrated by the Cremer−Kraka parameter H(r), at the M−C bond critical points are each independent of X and are constant for a given R in CR3. The values of ρ(r) are constant for a given M and ancillary ligand X and vary only slightly with varying M; the variation between R = H, F, and Cl is more pronounced with lower electron densities as R = H > F > Cl for each case. Similarly, the values of H(r) show completely analogous dependencies; the small negative values are indicative of strong ionic components to the M−C bonds in each case, with a greater ionic character when X = Cl > F > H, consistent with bond ionicities. The overall picture is consistently one of strongly polar M−CR3 bonds. For each metal, the order of M−CR3 bond length is R = Cl > F > H. This correlates inversely with the s character in the carbon orbital, indicating that any spatial contraction of this orbital is not dominant in determining bond lengths in these compounds. While the signs of the local charges on M and C in M−CH3 (δ+/δ−) and M−CF3 (δ+/δ+) are consistent with a longer M bond to CF3, they do not explain an even longer bond to CCl3. While local charges on carbon are constant for each M−CR3, significant variation in the magnitude of the positive charge on M generates no significant difference in the M−C distances, the electron density ρ(r), or Cremer−Kraka parameter H(r) at the bond critical points, or the bond ionicities. We considered that the longer bonds to CF3 and CCl3 might be a result of steric crowding in the somewhat constrained space in the “wedge” of the Cp2M system. It seemed sensible to design a model system in which steric effects would be minimized. The complexes in Tables 3 and 4 are pseudotetrahedral M(IV) complexes with d 0 electron counts and 16-valence electron configurations. The NBO method has been widely used to describe the bonding in transition-metal complexes, for which the valence orbital set is ns and (n − 1)d, with essentially no contributions from the energetically high-lying np orbitals.51c,54 As a consequence of a valence orbital set of six, the maximum number of electron pairs, bonding and nonbonding, that can be accommodated around a transition metal must also be six,

Figure 5. Top: one of a degenerate pair of Kohn−Sham LUMOs in Zr(CH3)4 (55) and the corresponding HOMO in Ru(CH3)4 (64). Bottom: the corresponding NBOs. Each orbital is depicted as a 0.05e isosurface.

Considering, first, the group 4 complexes and comparing data from Tables 4 and 5 in which two Cp ligands have been replaced by two CH3 groups, the results are remarkably consistent. Replacement of Cp by CH3 causes a dramatic increase in the NPA positive charge at the metal center, illustrating how superior a donor Cp is compared with CH3. The negative charge at the CH3 and CCl3 carbons increases and the positive charge at the CF3 carbon decreases slightly, with charges on the H, Cl, and F atoms remaining essentially constant. Bond ionicities are slightly higher compared with those of Cp2M analogues, particularly for Ti, but the values of ρ(r) and H(r) 1492

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a

1493

1.996

Os(CH3)3(CCl3) (69)

Delocalization per CH3 group.

2.041 2.002

Os(CH3)4 (67) Os(CH3)3(CF3) (68)

2.027 1.997

Ru(CH3)4 (64) Ru(CH3)3(CF3) (65)

1.977

1.891

Fe(CH3)3(CCl3) (63)

Ru(CH3)3(CCl3) (66)

2.078 2.184 2.151 2.246 2.352 2.314 2.220 2.323 2.285 1.891 1.905

Ti(CH3)4 (52) Ti(CH3)3(CF3) (53) Ti(CH3)3(CCl3) (54) Zr(CH3)4 (55) Zr(CH3)3(CF3) (56) Zr(CH3)3(CCl3) (57) Hf(CH3)4 (58) Hf(CH3)3(CF3) (59) Hf(CH3)3(CCl3) (60) Fe(CH3)4 (61) Fe(CH3)3(CF3) (62)

compound

M− CR3

2.042

2.041 2.035

2.031

2.027 2.021

1.893

2.078 2.056 2.063 2.246 2.221 2.227 2.220 2.197 2.203 1.891 1.881

M− CH3

1.831

1.370

1.814

1.367

1.802

1.364

1.374 1.810

1.373 1.811

1.370 1.805

C−R

1.098

1.100 1.099

1.095

1.097 1.096

1.094

1.097 1.096 1.096 1.098 1.098 1.097 1.099 1.098 1.098 1.096 1.096

C−H

distances (Å)

109.0

109.5 112.5

109.4

109.5 112.4

108.0

109.5 109.2 107.1 109.5 109.3 108.0 109.5 109.0 107.6 109.5 111.5

CH− M− CH

109.8

106.2

109.6

106.2

111.0

107.3

109.9 111.1

109.6 111.0

109.8 111.7

CH− M− CR

0.74

0.74 0.70

0.49

0.50 0.47

0.57

1.50 1.42 1.43 1.99 1.92 1.97 2.20 2.12 2.16 0.62 0.55

qM

angles (deg)

qCR

0.02

−0.29 −0.01

−0.76 0.21

0.01

0.88 −0.38

−0.25

0.93 −0.37

−0.31

0.89 −0.37

0.57 −0.37 −0.66 0.03

0.61 −0.37 −0.62 0.03

0.69 −0.37 −0.52 0.03

qCR

−0.77 0.20 −0.77 0.21

−0.68 0.20

−0.70 0.19 −0.69 0.20

−0.73 0.20

0.22 0.23 0.23 0.22 0.23 0.23 0.23 0.23 0.23 0.20 0.21

qCH

−1.04 −1.03 −1.02 −1.17 −1.16 −1.17 −1.23 −1.23 −1.23 −0.75 −0.73

qCH

atomic NPA charges

−0.22

−0.54 −0.57

−0.50 −0.53

−0.42 −0.43

CR3

−0.13 −0.32

−0.17 −0.14 −0.26

−0.08 −0.22

−0.13 −0.09 −0.18

−0.13 −0.25

−0.38 −0.34 −0.36 −0.51 −0.50 −0.51 −0.54 −0.54 −0.54 −0.15 −0.10

CH3

group charge

25%

20%

22%

18%

32%

22%

62% 73%

58% 69%

50% 59%

M− CR

14%

18% 15%

9%

13% 9%

13%

46% 44% 43% 56% 54% 56% 62% 60% 60% 18% 13%

M− CH

bond ionicities

sp3.7

sp3.6 sp3.8

sp4.3

sp4.1 sp4.3

sp4.6

sp3.0 sp3.2 sp3.0 sp2.7 sp2.8 sp2.7 sp2.5 sp2.6 sp2.5 sp4.3 sp4.7

CH3

sp2.0

sp2.0

sp2.2

sp2.1

sp2.4

sp2.2

sp1.6 sp1.5

sp1.7 sp1.6

sp sp1.7

1.8

CR3

C → M hybrid

Table 5. Metric, NBO, NPA, and QTAIM Data for Complexes M(CH3)3(CR3) (M = Ti, Zr, Hf, Fe, Ru, Os; R = H, F, Cl)

H(r) M− CR 0.109 0.115 0.114 0.094 0.099 0.098 0.099 0.105 0.104 0.131 0.134

ρ(r) M− CH

0.158 −0.066 0.137

0.137 0.149 −0.057 0.139

0.143 −0.050 0.129

0.129 0.138 −0.045 0.131

0.134 −0.059 0.131

0.131 −0.054

0.081 −0.018 0.087 −0.021

0.075 −0.014 0.081 −0.017

0.087 −0.026 0.093 −0.028

ρ(r) M− CR

−0.056

−0.055 −0.057

−0.048

−0.047 −0.050

−0.055

0

0 0

0

0 0

0

9a 10a 11a 9a 10a 11a 9a 10a 11a 0 0

0 0 1(CH)a 5(CF) 3(CH)a 5(CCl) 1(CH)a 5(CH)a 14(CF) 4(CH)a 15(CCl) 4(CH)a 8(CH)a 20(CF) 7(CH)a 27(CCl) 6(CH)a

0 0

0 0

Md → CRσ*

ΔEdelocalization (kcal/mol) CHσ → Md

−0.041 −0.046 −0.045 −0.026 −0.029 −0.029 −0.030 −0.033 −0.032 −0.055 −0.058

H(r) M− CH

QTAIM BCP data (M−C)

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remain very similar. Hybridizations for the carbon donor orbital in CR3 show higher p character in all cases, but the trend of significantly lower p character for R = F and Cl remains. The metal orbital hybrids are very close to the expected51c sd3 in all cases. We conclude that M(CH3)4 and M(CH3)3(CR3) complexes are good models for the Cp2M analogues and that the f undamental electronic characteristics of the M−CCR3 bonds are not signif icantly disturbed. Comparing M(CH3)4, M(CH3)3(CF3), and M(CH3)3(CCl3) illustrates that, in these less sterically hindered systems, the M−CCF3 bond is always the longest, with a shorter M−CCCl3 bond and an even shorter M−CCH3 distance. A CF3 or CCl3 group slightly shortens the M−CCH3 bond in M(CH3)3(CR3) compared with that in the corresponding M(CH3)4 analogue. Thus, the trend in M−C distances now follows the trend in relative local charges on the atoms in the M−C bonds, with M−CH3 being shorter than M−CCCl3, and with M−CCF3 being the longest, presumably due to δ+/δ+ repulsion. In these strongly polar bonds, the M−C distance does not correlate with the s character in the carbon donor orbital. Turning to the d4 group 8 analogues in Table 5, we note that all bond lengths decrease compared with those of the appropriate group 4 analogues, as expected for the smaller atomic radii. While the Fe−CCF3 distance remains slightly longer than the analogous Fe−CCH3 distances, the Fe−CCCl3 distance is now the same as Fe−CCH3. However, examination of the Ru and Os analogues reveals that the M−CCF3 and M−CCCl3 distances both become shorter than M−CCH3, with M−CCCl3 even becoming slightly shorter than M−CCF3! The trend for CH3 and CF3 is in line with other previously discussed examples from later in the periodic table. The NPA charges on the metals decrease substantially compared with their more electropositive group 4 counterparts; charges on the CH3 and CCl3 carbons become less negative and those on the CF3 carbons more positive, with those on H, Cl, and F remaining essentially unchanged from the values observed in group 4. Thus, once again, most of the changes in charge distribution are accommodated by M and C. The bond ionicities for the group 8 compounds are substantially lower than those for group 4; this additional covalency is also reflected in greater values for the electron densities ρ(r) and more negative values of the Cremer−Kraka parameter H(r) at the M−C bond critical points. In these more covalent compounds, there is no correlation between charges on the metals or ligated carbon atoms and the M−C bond length variations, but a more sensible correlation with increased s character in the carbon orbital. NBO analysis provides further insight and suggests that, for group 4, additional π-components serve to enhance the relative shortening of M−CCH3, whereas in group 8, π-components serve to shorten both M−CCF3 and M−CCCl3 relative to M−CCH3. As discussed above, the principal electronic structural difference between the groups 4 and 8 complexes involves the occupancies (vacant or filled, respectively) of the degenerate pair of d orbitals. NBO recognizes delocalizations from the reference structure and allows their depiction as natural localized molecular orbitals (NLMOs).51c Figure 6 shows the NLMOs of C−H σ-bonds in Zr(CH3)4 (55) and in Ru(CH3)4 (64). In complex 55, partial delocalization from the CH(σ)→Zr(4d) is apparent, whereas no such delocalization is observed in the analogous Ru complex (64); this donation from all 12 available CH(s) orbitals provides an overall occupancy of the two formally empty

Figure 6. NLMOs for a single C−H(σ) in Zr(CH3)4 (55) and in Ru(CH3)4 (64) represented as 0.03e isosurfaces.

Zr(4d) orbitals of 0.216e, or 0.054e per CH3 group. Figure 7 illustrates NLMOs of C−F(σ*) in Zr(CH3)3(CF3) (56) and Ru(CH3)3(CF3) (65), illustrating significant delocalization from the filled Ru(4d) → CF(σ*), but none, of course, from the empty Zr(4d). In this case, the two Ru(4d) orbitals are

Figure 7. NLMOs for a single C−F(σ*) in Zr(CH3)3(CF3) (56) and Ru(CH3)3(CF3) (65) represented as 0.03e isosurfaces.

depleted by 0.116e. Figure 8 shows the NLMOs of one of the Zr(4d) and Ru(4d) orbitals; clearly, the Zr(4d) is populated exclusively by delocalization from CH(σ) orbitals with negligible contributions from F lone pairs, whereas the Ru(4d)

Figure 8. NLMOs for Zr(4d) and Ru(4d) in Zr(CH3)3(CF3) (56) and Ru(CH3)3(CF3) (65) represented as 0.03e isosurfaces.

is depopulated almost exclusively by delocalization into CF(σ*) orbitals. It is well established56 that such delocalizations can be quantified using second-order perturbative NBO analysis for estimating metal−ligand π-bonding interaction energies (ΔE), obtained from the off-diagonal Fock matrix element expressed in the NBO basis.51c Values of ΔE reflect stabilization energies resulting from delocalization from the reference structure, and previous studies have generally involved delocalization of metalbased lone pairs into π* antibonding orbitals of the ligands. Delocalization energies for the groups 4 and 8 model compounds are presented in Table 5; for each group 4 complex, the π-donation is 9−10 kcal/mol/methyl group for Ti, Zr, and Hf, whereas for the group 8 analogues, the π-acceptor properties of the CF3 and CCl3 ligands dominate over the relatively small 1494

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qFγ

−0.33 −0.33 −0.33 −0.33 −0.33 −0.33

qFβ

−0.35 −0.35 −0.34 −0.35 −0.35 −0.35

qFα

−0.36 −0.37 −0.37 −0.37 −0.37 −0.38

qCF3

1.00 1.00 1.00 1.00 1.00 1.00 0.63 0.62 0.63 0.63 0.63 0.63

qCγ qCβ

0.62 0.62 0.62 0.62 0.62 0.63 0.35 0.27 0.26 0.55 0.60 0.55

qCα qM

1.45 1.95 2.15 0.58 0.48 0.72 108.6 108.6 108.6 108.6 108.6 108.6

CF3 γ-CF2

108.7 108.8 108.8 108.7 108.7 108.8 107.6 107.4 107.3 107.7 107.6 107.4

β-CF2

1.351 1.352 1.351 1.351 1.351 1.351

α-CF2 CF3

1.339 1.339 1.339 1.339 1.339 1.339

γ-C−F

106.5 106.2 106.2 106.8 106.4 105.7

CONCLUSIONS The first series of titanocene trifluoromethyl halide and pseudohalide complexes, Cp2Ti(CF 3)(X) (where X = F (1), OTf (2), Cl (12), Br (13), I (14), N3 (15), and OSO2Ph (16)), has been prepared. Our approach capitalized on the latent reactivity of titanium fluoride bonds; reaction between Cp2TiF2 and (CH3)3SiCF3 provided access to Cp2Ti(CF3)(F) (1), which served as a platform for preparing other Cp2Ti(CF3)(X) complexes using similar (CH3)3Si−F elimination chemistry. This collection of compounds represents a milestone in the history of organometallic fluorocarbon chemistry. In contrast to earlier beliefs, the Ti−CF3 linkage in these complexes is remarkably robust, showing no evidence of an α-fluoride (Ti···F−CF2) between the electrophilic Ti(IV) metal center and any of the C−F bonds in the trifluoromethyl group. However, in the solid-state, these complexes are shock-sensitive, and the

1.364 1.366 1.366 1.361 1.361 1.362

β-C−F

■

1.385 1.388 1.390 1.378 1.381 1.390

α-C−F M−CF

2.201 2.329 2.337 1.930 2.017 2.015

M

Ti (70) Zr (71) Hf (72) Fe (73) Ru (74) Os (75)

M−CH

π-acceptance by each CH3 group, with CCl3 being as good a πacceptor as CF3. The π-donor properties of the metal increase significantly from Fe→Ru→Os. All things considered, we conclude that, for group 4 complexes, π-bonding is a significant factor in shortening the strongly ionic M−CCH3 relative to M−CCR3 (R = F, Cl), for the predominantly covalent group 8 analogues, π-back-bonding helps to shorten the predominantly covalent M− to M−CCR3 relative to M−CCH3, at least for Ru and Os. Despite the presence or absence of π-donation into the CF(σ*), the C−F bond lengths remain identical to within 0.01 Å in all compounds. In fact, the Ru and Os compounds in which this back-bonding is at its greatest actually have slightly shorter C−F bonds than the group 4 analogues, for which no back-bonding exists. It is not clear what is “normal” for a C−F bond length in such compounds, and comparison between different molecules usually introduces multiple variables. Accordingly, the molecules M(CH3)3(CF2CF2CF2CF3) (M = Ti (70), Zr (71), Hf (73), Fe (74), Ru (75), Os (76)) were studied computationally, using the same DFT methodology, and results are presented in Table 6. An internal comparison between a C−F bond α, β, and γ to a metal can now be made; this is a more informative intramolecular comparison and avoids comparisons between CF3 groups, which necessarily must be intermolecular. The trends in M−CCH3 versus M−CCF2 distances are identical to those in Table 5, with Ru and Os having shorter bonds to the fluoroalkyl carbon than to CH3. The same relative trends are also observed for the α-C−F bond distances when compared to the α-C−F distances in Table 5, though it is noteworthy that the C−F bonds α to the metals are all significantly shorter than the C−F bonds in the CF3 group at the end of the fluoroalkyl chain (Table 6). In the internal comparison set, the α-CF2 distances are the longest and most variable, with significant shortening observed in the β-CF2 groups; on reaching the γ-CF2 groups, convergence to a constant value of 1.351 Å is achieved. Application of Bent’s Rule40a predicts that replacement of a fluorine in a CF3 group with anything less electronegative than F (i.e., anything) should result in less competition for p character between the remaining fluorines and hence longer CF bonds; comparison between the γ-CF2 and the terminal CF3 distances illustrates this nicely. We suggest that the significant elongation of C−F bonds α to metals is mainly a consequence of the electropositivity of the metals, with minor, if any, contributions from π-effects; the bond-lengthening effect is most pronounced at the α-position and decays rapidly on moving away from the metal.

2.055 2.220 2.195 1.882 2.021 2.034

NPA charges angles (deg) distances (Å)

Table 6. Selected Metric and NPA Data for Complexes M(CH3)3(CF2CF2CF2F3) (M = Ti, Zr, Hf, Fe, Ru, Os)

Article

−0.33 −0.33 −0.33 −0.33 −0.33 −0.33

qCF3

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(Note: The reaction also works with only THF as the solvent, but slightly higher yields of 1 are achieved in a 1:1 mixture of THF and Et2O). While stirring, Me3SiCF3 (1.0 mL, 0.99 g, 6.94 mmol) was added dropwise, and the resulting yellow/brown slurry was stirred at room temperature for 15 h. The mixture was then filtered through a glass-fiber plug to remove insoluble materials. The orange/brown filtrate was collected and the volatile materials were removed under reduced pressure to produce an orange/brown waxy solid. (Note: Care was taken to remove the source of vacuum pressure as soon as a waxy solid was formed, as complete removal of solvent causes rapid decomposition of the compound.) The solid was washed with cold hexanes (3 × 3 mL), and any residual volatiles were removed under reduced pressure (again, by exposing the mixture to reduced pressure for a minimum amount of time to avoid decomposition) to produce 1 as a bright yellow solid (110 mg, 0.413 mmol, 60%). Complex 1 decomposes at room temperature and must be stored cold (−30 °C). Safety Reminder! The quantities in which 1 is synthesized should be limited to ∼150 mg, as 1 is susceptible to violent decomposition and must be handled with care. 1H NMR (300 MHz, THF-d8, RT): δ 6.40 (d, 3JF−H = 1.0 Hz, Cp). 19F NMR (282 MHz, THF-d8, RT): δ 206.3 (s, Ti-F), −24.0 (s, Ti-CF3). 13C{1H} NMR (75 MHz, THF-d8, RT): δ 153.1 (dq, 1JC−F = 382.6 Hz, 2JC−F = 4.8 Hz, Ti-CF3), 117.4 (s, Cp). IR (cm−1, Nujol mull): 585 (m), 620 (w), 692 (w), 822 (s), 836 (s), 849 (s), 941 (m), 979 (s), 992 (s), 1017 (w), 1029 (w), 1069 (s), 1081 (s), 1133 (w), 1201 (w), 1274 (w), 1440 (m). Synthesis of Cp2Ti(CF3)(OTf) (2). In the drybox, a 20 mL scintillation vial equipped with a stir bar was charged with 1 (100 mg, 0.376 mmol). The minimum volume of toluene (∼5 mL) required to completely dissolve 1 was added. While stirring, Me3SiOTf (68 μL, 84 mg, 0.376 mmol) was added as one aliquot, and pyridine (36 μL, 36 mg, 0.451 mmol) added immediately thereafter. The resulting orange solution was layered with hexanes (∼2 mL) and placed in a freezer (−30 °C) to crystallize. After 15 h, orange-red crystals were observed, and the solvent was removed by pipet. The crystals were washed with cold hexanes (∼5 mL), and residual hexanes were removed by brief exposure to reduced pressure. To remove traces of 1 still present in the product mixture, these crystals were allowed to sit at RT for 2 h and then dissolved in a pyridine (36 μL, 36 mg, 0.451 mmol)/toluene (∼5 mL) solution. Hexanes were added (∼2 mL), and the mixture was placed again in the freezer to crystallize. After 15 h, orange-red crystals were formed, and the solvent was removed by pipet. After a wash with cold hexanes (∼5 mL), the crystals were exposed briefly to reduced pressure to remove residual hexanes, and 2 (105 mg, 0.265 mmol, 70%) was isolated. Complex 2 decomposes at room temperature within a few hours and should be stored cold (−30 °C). Safety Reminder! The quantities in which this compound is synthesized should be limited to ∼150 mg, as 2 is susceptible to violent decomposition and must be handled with care. 1H NMR (300 MHz, THF-d8, 0 °C): δ 6.76 (s, Cp). 19F NMR (282 MHz, THF-d8, 0 °C): δ −25.6 (s, Ti-CF3), −80.0 (s, OTf). 13C{1H} NMR (75 MHz, THF-d8, 0 °C): δ 151.8 (q, 1JC−F = 382.9 Hz, Ti-CF3), 120.7 (s, Cp), 120.2 (q, 1JC−F = 317.3 Hz, OTf). IR (cm−1, Nujol mull): 594 (w), 632 (m), 694 (w), 766 (w), 836 (s), 857 (w), 864 (w), 951 (sh, w), 990 (s), 1015 (s), 1026 (sh, w), 1069 (s), 1082 (s), 1132 (w), 1175 (s), 1184 (sh, m), 1204 (s), 1237 (s), 1298 (w), 1328 (w), 1345 (s). Synthesis of Cp2Ti(CF3)(Cl) (12). Method A. In the drybox, a 20 mL scintillation vial equipped with a stir bar was charged with 1 (100 mg, 0.376 mmol) and 5 mL of THF. While stirring, (CH3)3SiCl (0.24 mL, 1.88 mmol) was added by syringe. The reaction mixture was allowed to stir for 3 h at room temperature, and then all volatile materials were removed under reduced pressure to produce an orange solid. This solid was dissolved in toluene (3−4 mL), layered with hexanes (1−2 mL), and placed in a freezer (−30 °C) to crystallize. After 15 h, orange crystals were observed, and the solvent was removed by pipet. The crystals were washed with cold hexanes (3 × 1 mL) and the residual hexanes were removed by brief exposure to reduced pressure to give complex 12 (93 mg, 0.329 mmol, 88%). Complex 12 decomposes within a day at room temperature and should be stored cold (−30 °C). Safety Reminder! The quantities in which 1 is

energetic decomposition of Cp2Ti(CF3)(F) (1) provides a new entry into fluorinated materials containing −(CF2−CF2)− and −(CF2−CFH)− units. These Ti−CF3 complexes also provide an unprecedented opportunity to understand the electronic structure and bonding of early, middle, and late transition-metal trifluoromethyl, methyl, fluoroalkyl, and alkyl complexes. In contrast to middle- and laterow transition metals, the Ti−CF3 complexes featured a longer titanium−carbon bond compared with their Ti−CH3 counterparts. Density functional theory (DFT) calculations were performed to gain insight into this unusual bonding scenario. The calculations accurately predicted longer Ti−CF3 distances than for each specific CH3 analogue, and the trend extended to structurally related Zr and Hf analogues. Simpler model compounds from groups 4 and 8 showed that, for group 4 complexes, π-bonding is a significant factor in shortening the strongly ionic M−CH3 relative to M−CF3, whereas for the group 8 analogues, π-back-bonding helps to shorten the predominantly covalent M−CF3 relative to M−CH3. The bonding analysis suggests that the significant elongation of C−F bonds α to metals is mainly a consequence of the electropositivity of the group 4 metal centers, with minor, if any, contributions from π-effects; the bond-lengthening effect is most pronounced at the α-position and decays rapidly on moving away from the metal.

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EXPERIMENTAL SECTION

General Considerations. All reactions and manipulations were carried out using a Vacuum Atmospheres (MO 40-2 Dri-train) inert atmosphere (He) drybox, or using standard Schlenk techniques. Glassware was dried overnight at 150 °C before use. All solution NMR spectra were obtained in C6D6 using a Bruker Avance 300 MHz spectrometer. Chemical shifts for 1H and 13C{1H} NMR spectra were referenced to solvent impurities. For 19F NMR spectra, CFCl3 was used as an external reference at δ 0.00 ppm. All IR data were recorded on a Thermo-Nicolet FT-IR spectrometer. The SEM experiments were performed using a JEOL JSM-6460 scanning electron micrsocope. The powder X-ray diffraction experiments were performed using a Bruker D8 Discover with a GADDS diffractometer. The bonding configuration and the compositional analysis of the fluorocarbon materials were carried out by X-ray photoelectron spectroscopy (SPECS, Germany) using a hemispherical analyzer (HSA 3500). All solid-state NMR spectra were obtained using a Bruker Avance 400 MHz spectrometer with a wide-bore superconducting magnet. Samples were spun at ∼20 000−30 000 Hz in a 2.5 mm zirconia rotor at the magic angle. The 1H, 19F, and 13C frequencies were 399.95, 376.32, and 100.58 MHz, respectively. The 13C spectra were collected under cross-polarization from 1H. The number of acquisitions for the 13C cross-polarization magic-angle spinning (CPMAS) experiment was ∼20 000, and 800 data acquisitions were needed for the 19F MAS results. Materials. Unless otherwise noted, reagents were purchased from commercial suppliers and used without further purification. THF-d8 (Cambridge Isotope Laboratories) was dried over a Na mirror prior to use. Celite (Aldrich), 4 Å molecular sieves (Aldrich), and alumina (Brockman I, Aldrich) were dried under dynamic vacuum at 250 °C for 48 h prior to use. All solvents (Aldrich) were purchased anhydrous and were dried over KH for 24 h, passed through a column of activated alumina, and stored over activated 4 Å molecular sieves prior to use. Complexes 1 and 2 were prepared according to literature procedures.11 Caution! The quantities in which Cp2Ti(CF3)(F) (1), Cp2Ti(CF3)(OTf) (2), and Cp2Ti (CF3)(Cl) (12) are synthesized should be limited to ∼150 mg, as 1, 2, and 12 are susceptible to violent decomposition and must be handled with care. Synthesis of Cp2Ti(CF3)(F) (1). In the drybox, a 20 mL scintillation vial equipped with a stir bar was charged with Cp2TiF2 (150 mg, 0.694 mmol) and CsF (105 mg, 0.694 mmol). Tetrahydrofuran (5 mL) and Et2O (5 mL) were added to create a suspension 1496

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synthesized should be limited to ∼150 mg, as 1 is susceptible to violent decomposition and must be handled with care. Method B. A 20 mL scintillation vial equipped with a stir bar was charged with 1 (100 mg, 0.376 mmol) and toluene (5 mL). While stirring, (CH3)3SiCl (0.24 mL, 1.88 mmol) was added by syringe. The reaction mixture was allowed to stir for 3 h at room temperature, layered with hexanes (1−2 mL), and placed in a freezer (−30 °C) to crystallize. After 24 h, X-ray quality orange crystals of 12 formed. 1H NMR (300 MHz, THF-d8, RT): δ 6.52 (s, Cp). 19F NMR (282 MHz, THF-d8, RT): δ −21.2 (s, Ti-CF3). 13C{1H} NMR (75 MHz, THF-d8, RT): δ 153.2 (q, 1JC−F = 382.1 Hz, Ti-CF3), 118.6 (s, Cp). IR (cm−1, Nujol mull): 691 (w), 826 (s), 843 (w), 854 (w), 873 (w), 933 (m), 976 (s), 1015 (m), 1026 (w), 1064 (s), 1077 (s). Synthesis of Cp2Ti(CF3)(Br) (13). In the drybox, a 20 mL scintillation vial equipped with a stir bar was charged with 1 (100 mg, 0.376 mmol) and THF (5 mL). While stirring, (CH3)3SiBr (51 μL, 0.395 mmol) was added by syringe. The reaction mixture was allowed to stir for 45 min at room temperature, and then all volatile materials were removed under reduced pressure to give an orange-red solid. This solid was dissolved in toluene (3−4 mL), layered with hexanes (1−2 mL), and placed in a freezer (−30 °C) to crystallize. After 15 h, orange-red crystals formed, and the solvent was removed by pipet. The crystals were washed with cold hexanes (3 × 1 mL); residual hexanes were removed by brief exposure to reduced pressure to produce complex 13 (103 mg, 0.315 mmol, 84%). Complex 13 decomposes within a few hours at room temperature and should be stored cold (−30 °C). 1 H NMR (300 MHz, THF-d8, RT): δ 6.58 (s, Cp). 19F NMR (282 MHz, THF-d8, RT): δ −20.6 (s, Ti-CF3). 13C{1H} NMR (75 MHz, THF-d8, RT): δ 155.2 (q, 1JC−F = 381.5 Hz, Ti-CF3), 118.6 (s, Cp). IR (cm−1, Nujol mull): 690 (w), 827 (s), 844 (w), 854 (w), 871 (w), 932 (m), 975 (s), 1015 (m), 1024 (w), 1064 (s), 1074 (s). Synthesis of Cp 2Ti(CF 3)(I) (14). In the drybox, a 20 mL scintillation vial equipped with a stir bar was charged with 1 (93 mg, 0.351 mmol) and toluene (5 mL). While stirring, (CH3)3SiI (50 μL, 0.351 mmol) was added as one aliquot, and pyridine (34 μL, 0.421 mmol) immediately added. (Note: Pyridine was added to help stabilize the product that is formed. Without the presence of a Lewis base, rapid decomposition occurred in solution to form intractable solids.) The resulting green solution was layered with hexanes (∼2 mL) and placed in a freezer (−30 °C) to crystallize. After 15 h, a fine green-black solid was observed, and the solvent was removed by pipet. The solid was washed with cold hexanes (3 × 1 mL); residual hexanes were removed by brief exposure to reduced pressure to produce complex 14 (88 mg, 0.235 mmol, 67%). Complex 14 decomposes within a few hours at room temperature and should be stored cold (−30 °C). 1H NMR (300 MHz, THF-d8, −20 °C): δ 6.72 (s, Cp). 19F NMR (282 MHz, THF-d8, −20 °C): δ −20.5 (s, Ti-CF3). 13C{1H} NMR (75 MHz, THF-d8, −20 °C): δ 159.3 (q, 1JC−F = 380.0 Hz, Ti-CF3), 118.3 (s, Cp). IR (cm−1, Nujol mull): 689 (m), 822 (s), 857 (s), 869 (sh, w), 927 (w), 943 (m), 971 (s), 992 (m), 1015 (m), 1024 (m), 1061 (s), 1076 (sh, m), 1132 (m), 1202 (w). Synthesis of Cp 2Ti(CF 3)(N 3) (15). In the drybox, a 20 mL scintillation vial equipped with a stir bar was charged with 1 (100 mg, 0.376 mmol) and THF (5 mL). While stirring, (CH3)3SiN3 (0.49 mL, 3.76 mmol) was added by syringe. The reaction mixture was allowed to stir for 2 h at room temperature, and then all volatile materials were removed under reduced pressure to produce an orange-red solid. This solid was dissolved in toluene (3−4 mL), layered with hexanes (1−2 mL), and placed in a freezer (−30 °C) to crystallize. After 24 h, X-ray quality, orange-red crystals formed, and the solvent was removed by pipet. The crystals were washed with cold hexanes (3 × 1 mL); residual hexanes were removed by brief exposure to reduced pressure to produce complex 15 (77 mg, 0.266 mmol, 71%). Complex 15 decomposes within a few days at room temperature and should be stored cold (−30 °C). 1H NMR (300 MHz, THF-d8, −20 °C): δ 6.43 (s, Cp). 19F NMR (282 MHz, THF-d8, −20 °C): δ −23.2 (s, Ti-CF3). 13 C{1H} NMR (75 MHz, THF-d8, −20 °C): δ 153.9 (q, 1JC−F = 381.8 Hz, Ti-CF3), 116.8 (s, Cp). IR (cm−1, Nujol mull): 618 (w), 689 (m), 827 (s), 856 (sh, w), 938 (s), 972 (s), 1018 (m), 1027 (w), 1065 (s),

1131 (m), 1170 (w), 1201 (w), 1340 (s), 1964 (w), 2060 (s), 2663 (w), 2724 (w). Synthesis of Cp2Ti(CF3)(OSO2Ph) (16). In the drybox, a 20 mL scintillation vial equipped with a stir bar was charged with 1 (66 mg, 0.247 mmol) and toluene (3 mL). While stirring, (CH3)3Si(OSO2Ph) (50 μL, 0.247 mmol) was added as one aliquot, and pyridine (24 μL, 0.296 mmol) immediately added. (Note: Pyridine was added to help stabilize the product that is formed. Without the presence of a Lewis base, rapid decomposition occurred in solution to form intractable solids.) The resulting orange solution was layered with hexanes (∼1 mL) and placed in a freezer (−30 °C) to crystallize. After 15 h, a fine orange solid was observed, and the solvent was removed by pipet. The solid was washed with cold hexanes (3 × 1 mL); residual hexanes were removed by brief exposure to reduced pressure to produce complex 16 (73 mg, 0.181 mmol, 73%). Complex 16 decomposes within a few hours at room temperature and should be stored cold (−30 °C). 1H NMR (300 MHz, THF-d8, RT): δ 7.80 (m, Ph), 7.48 (m, Ph), 7.45 (m, Ph), 6.63 (s, Cp). 19F NMR (282 MHz, THF-d8, RT): δ −24.5 (s, Ti-CF3). 13C{1H} NMR (75 MHz, THF-d8, RT): δ 152.4 (q, 1JC−F = 383.1 Hz, Ti-CF3), 143.1 (s, Ph), 132.3 (s, Ph), 129.4 (s, Ph), 127.3 (s, Ph), 119.3 (s, Cp). IR (cm−1, Nujol mull): 582 (w), 612 (m), 691 (s), 727 (s), 763 (m), 832 (s), 844 (m), 855 (sh, w), 871 (sh, w), 940 (w), 974 (s), 1004 (w), 1028 (w), 1080 (s), 1110 (m), 1128 (w), 1168 (s), 1208 (w), 1223 (w), 1239 (w), 1302 (s). X-ray Crystallographic Details for Cp2Ti(CF3)(Cl) (12). A crystal (0.24 × 0.20 × 0.14 mm3) of complex 12 was mounted in a nylon cryoloop using Paratone-N oil under an argon gas flow. The data were collected on a Bruker Platform diffractometer with a 1K CCD and cooled using a Bruker KRYO-FLEX liquid nitrogen vapor cooling device. The instrument was equipped with a sealed, graphite monochromatized Mo Kα X-ray source (λ = 0.71073 Å) and a 0.5 mm monocapillary. A hemisphere of data was collected using φ scans, with 30 s frame exposures and 0.3° frame widths. Data collection and initial indexing and cell refinement were handled using SMART software.57 Frame integration, including Lorentz-polarization corrections, and final cell parameter calculations were carried out using SAINT software.58 The data were corrected for absorption using the SADABS program.59 Decay of reflection intensity was monitored by analysis of redundant frames. The structure was solved using direct methods and difference Fourier techniques. A substitutional chloride impurity was modeled at the CF3 site, with site occupancy factors tied to one. The site occupancy factor on the chloride impurity refined to 0.14(1). A racemic twin was also refined, and the batch scale factors converged to 0.54(7). Nonhydrogen atoms were refined anisotropically, and hydrogen atoms were treated as idealized contributions. Structure solution, refinement, graphics, and creation of publication materials were performed using SHELXTL.60 Additional details regarding data collection are provided in the CIF file (Supporting Information). X-ray Crystallographic Details for Cp 2Ti(CF3)(N3) (15). A crystal (0.30 × 0.10 × 0.10 mm3) of complex 15 was mounted in a nylon cryoloop using Paratone-N oil under an argon gas flow. The data were collected on a Bruker D8 APEX II charge-coupled device (CCD) diffractometer with an American Cryoindustries Cryocool lowtemperature device. The instrument was equipped with a graphite monochromatized Mo Kα X-ray source (λ = 0.71073 Å) and a 0.5 mm monocapillary. A hemisphere of data was collected using ω scans, with 10 s frame exposures and 0.5° frame widths. Data collection and initial indexing and cell refinement were handled using APEX II software.61 Frame integration, including Lorentz-polarization corrections, and final cell parameter calculations were carried out using SAINT+ software.62 The data were corrected for absorption using redundant reflections and the SADABS program.59 Decay of reflection intensity was not observed as monitored by analysis of redundant frames. The structure was solved using direct methods and difference Fourier techniques. Non-hydrogen atoms were refined anisotropically, and hydrogen atoms were treated as idealized contributions. Structure solution, refinement, graphics, and creation of publication materials were performed using SHELXTL.60 Additional details regarding data collection are provided in the CIF file (Supporting Information). 1497

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Computational Methods. All reported DFT calculations were carried out, without any symmetry constraints, using the robust hybrid B3LYP functional43 and the triple-ζ LACV3P**++ basis set,44 as implemented in the Jaguar63 suite of programs. Vibrational frequencies were calculated using analytic second derivatives, and all structures were confirmed as minima by the absence of imaginary frequencies. NBO/NPA calculations were done using the B3LYP-optimized structures, as part of the Jaguar suite. QTAIM calculations were run using the AIMAll program.64

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Rheingold, A. L. J. Am. Chem. Soc. 2001, 123, 7443−7444. (g) Bowden, A. A.; Hughes, R. P.; Lindner, D. C.; Incarvito, C. D.; Liable-Sands, L. M.; Rheingold, A. L. J. Chem. Soc., Dalton Trans. 2002, 3245−3252. (h) Hughes, R. P.; Ward, A. J.; Golen, J. A.; Incarvito, C. D.; Rheingold, A. L.; Zakharov, L. N. Dalton Trans. 2004, 2720−2727. (i) Bourgeois, C. J.; Hughes, R. P.; Husebo, T. L.; Smith, J. M.; Guzei, I. M.; Liable-Sands, L. M.; Zakharov, L. N.; Rheingold, A. L. Organometallics 2005, 24, 6431−6439. (j) Hughes, R. P.; Maddock, S. M.; Guzei, I. A.; Liable-Sands, L. M.; Rheingold, A. L. J. Am. Chem. Soc. 2001, 123, 3279−3288. (7) (a) Klabunde, K. J.; Key, M. S.; Low, J. Y. F. J. Am. Chem. Soc. 1972, 94, 999−1000. (b) Lagow, R. J.; Gerchman, L. L.; Jacob, R. A.; Morrison, J. A. J. Am. Chem. Soc. 1975, 97, 518−522. (8) (a) Morrison, J. A.; Gerchman, L. L.; Eujen, R.; Lagow, R. J. J. Fluorine Chem. 1977, 10, 333−339. (b) Krause, L. J.; Morrison, J. A. J. Am. Chem. Soc. 1981, 103, 2995−3001. (9) Berthold, H. J.; Groh, G. Angew. Chem., Int. Ed. 1963, 2, 398−399. (10) (a) Ontiveros, C. D. Ph.D. Thesis, Unversity of Chicago, Chicago, IL, 1986. (b) Jones, W. D. Dalton Trans. 2003, 3991−3995. (11) Taw, F. L.; Scott, B. L.; Kiplinger, J. L. J. Am. Chem. Soc. 2003, 125, 14712−14713. (12) (a) Wilkinson, G.; Birmingham, J. M. J. Am. Chem. Soc. 1954, 76, 4281−4284. (b) Druce, P. M.; Kingston, B. M.; Lappert, M. F.; Spalding, T. R.; Srivastava, R. C. J. Chem. Soc. A 1969, 14, 2106−2110. (13) (a) Prakash, G. K. S.; Yudin, A. K. Chem. Rev. 1997, 97, 757− 786. (b) Singh, R. P.; Shreeve, J. M. Tetrahedron 2000, 56, 7613−7632. (c) Prakash, G. K. S.; Hu, J. ACS Symp. Ser. 2005, 911, 16−56. (14) (a) Huang, D.; Caulton, K. G. J. Am. Chem. Soc. 1997, 119, 3185−3186. (b) Huang, D.; Koren, P. R.; Folting, K.; Davidson, E. R.; Caulton, K. G. J. Am. Chem. Soc. 2000, 122, 8916−8931. (c) Vicente, J.; Gil-Rubio, J.; Bautista, D. Inorg. Chem. 2001, 40, 2636−2637. (15) Dubinina, G. G.; Brennessel, W. W.; Miller, J. L.; Vicic, D. A. Organometallics 2008, 27, 3933−3938. (16) (a) In situ: Kim, J.; Shreeve, J. M. Org. Biomol. Chem. 2004, 2, 2728−2734. (b) Isolated: Dubinina, G. G.; Furutachi, H.; Vicic, D. A. J. Am. Chem. Soc. 2008, 130, 8600−8601. (17) (a) Vicente, J.; Gil-Rubio, J.; Guerrero-Leal, J.; Bautista, D. Organometallics 2004, 23, 4871−4881. (b) Vicente, J.; Gil-Rubio, J.; Guerrero-Leal, J.; Bautista, D. Organometallics 2005, 24, 5634−5643. (18) Grushin, V. V.; Marshall, W. J. J. Am. Chem. Soc. 2006, 128, 12644−12645. (19) (a) Naumann, D.; Kirij, N. V.; Maggiarosa, N.; Tyrra, W.; Yagupolskii, Y. L.; Wickleder, M. S. Z. Anorg. Allg. Chem. 2004, 630, 746−751. (b) Menjon, B.; Martinez-Salvador, S.; Gomez-Saso, M. A.; Fornies, J.; Falvello, L. R.; Martin, A.; Tsipis, A. Chem. Eur. J. 2009, 15, 6371−6382. (20) Zopes, D.; Kremer, S.; Scherer, H.; Belkoura, L.; Pantenburg, I.; Tyrra, W.; Mathur, S. Eur. J. Inorg. Chem. 2011, 273−280. (21) Werkema, E. L.; Messines, E.; Perrin, L.; Maron, L.; Eisenstein, O.; Andersen, R. A. J. Am. Chem. Soc. 2005, 127, 7781−7795. (22) For examples, see: (a) [Cp*2Ti(CH3)(THF)]+: Bochmann, M.; Jaggar, A. J.; Wilson, L. M.; Hursthouse, M. B.; Motevalli, M. Polyhedron 1989, 8, 1838−1843. (b) [CpTi(CH3)2]2(C10H8): Woehrle, T.; Thewalt, U. J. Organomet. Chem. 1993, 456, C21−C23. (c) Cp2Ti(CH3)2: Thewalt, U.; Woehrle, T. J. Organomet. Chem. 1994, 464, C17−C19. (23) Koola, J. D.; Roddick, D. M. Organometallics 1991, 10, 591− 597. (24) Rogers, R. D.; Atwood, J. L.; Rausch, M. D.; Macomber, D. W. J. Crystallogr. Spectrosc. Res. 1990, 20, 555−560. (25) 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. (26) (a) Doherty, N. M.; Hoffmann, N. W. Chem. Rev. 1991, 91, 553−573. (b) Caulton, K. G. New J. Chem. 1994, 18, 25−41. (27) Cotton, F. A.; McCleverty, J. A. J. Organomet. Chem. 1965, 4, 490. (28) Clark, G. R.; Hoskins, S. V.; Roper, W. R. J. Organomet. Chem. 1982, 234, C9−C12.

ASSOCIATED CONTENT

* Supporting Information S

Full details of crystallographic data for complexes 12 and 15. Optimized geometry coordinates and final ZPE-corrected energies for all computed structures. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected] (R.P.H.), [email protected] (J.L.K.). Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS For financial support, we acknowledge Los Alamos National Laboratory (Director’s PD Fellowships to F.L.T., A.E.C., A.H.M., and T.C.) and the LANL Laboratory Directed Research & Development program (J.L.K.). R.P.H. acknowledges the National Science Foundation for generous support, and Professor Clark Landis for useful suggestions. Finally, we thank Drs. Brady J. Gibbons and George J. Havrilla (both LANL) for their assistance with the SEM, XRD, and XRF measurements, and Dr. Marisa J. Monreal for her assistance with generating the SEM images.

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REFERENCES

(1) Drobny, J. G. Technology of Fluoropolymers; CRC Press LLC: Boca Raton, FL, 2001. (2) (a) Kiplinger, J. L.; Richmond, T. G.; Osterberg, C. E. Chem. Rev. 1994, 94, 373−431. (b) Richmond, T. G. Top. Organomet. Chem. 1999, 3, 243−269. (3) Kaminsky, W.; Sinn, H. Transition Metals and Organometallics as Catalysts for Olefin Polymerization; Springer-Verlag: New York, 1988. (4) (a) Coffield, T. H.; Kozikowski, J.; Closson, R. D. Abstracts of the Proceedings of the 5th International Conference on Coordination Chemistry; Chemical Society: London, England, 1959; p 126. (b) King, R. B. Organometallic Syntheses; Academic Press: New York, 1965; Vol. 1, pp 148−150. (5) For general reviews, see: (a) Treichel, P. M.; Stone, F. G. A. Adv. Organomet. Chem. 1964, 1, 143−220. (b) Bruce, M. I.; Stone, F. G. A. Prep. Inorg. React. 1968, 4, 177−235. (c) Morrison, J. A. Adv. Inorg. Chem. Radiochem. 1983, 27, 293−316. (d) Brothers, P. J.; Roper, W. R. Chem. Rev. 1988, 88, 1293−1326. (e) Hughes, R. P. Adv. Organomet. Chem. 1990, 31, 183−267. (f) Morrison, J. A. Adv. Organomet. Chem. 1993, 35, 211−239. (6) For examples of CF3I oxidative addition to d8 and d10 metals, see: (a) King, R. B.; Treichel, P. M.; Stone, F. G. A. J. Am. Chem. Soc. 1961, 83, 3593−3597. (b) Clark, H. C.; Manzer, L. E. J. Organomet. Chem. 1973, 59, 411−428. (c) Appleton, T. G.; Berry, R. D.; Hall, J. R.; Neale, D. W. J. Organomet. Chem. 1988, 342, 399−422. (d) Hughes, R. P.; Overby, J. S.; Williamson, A.; Lam, K.-C.; Concolino, T. E.; Rheingold, A. L. Organometallics 2000, 19, 5190−5201. (e) Hughes, R. P.; Sweetser, J. T.; Tawa, M. D.; Williamson, A.; Incarvito, C. D.; Rhatigan, B.; Rheingold, A. L.; Rossi, G. Organometallics 2001, 20, 3800−3810. (f) Hughes, R. P.; Williamson, A.; Sommer, R. D.; 1498

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not be determined due to significant disorder between the CH3 and the Cl ligands. (46) Hunter, W. E.; Hrncir, D. C.; Bynum, R. V.; Penttila, R. A.; Atwood, J. L. Organometallics 1983, 2, 750−755. (47) (a) Bader, R. F. W. Atoms In Molecules: A Quantum Theory; Oxford University Press: Oxford, England, 1994. (b) Matta, C. F., Boyd, R. J., Eds. The Quantum Theory of Atoms in Molecules; WileyVCH: Weinheim, Germany, 2007. (48) (a) Bader, R. F. W.; Gillespie, R. J.; MacDougall, P. J. Mol. Struct. Energ. 1989, 11, 1−51. (b) Popelier, P. L. A. Coord. Chem. Rev. 2000, 197, 169−189. (49) (a) Henn, J.; Ilge, D.; Leusser, D.; Stalke, D.; Engels, B. J. Phys. Chem. A 2004, 108, 9442−9452. (b) Kovács, A.; Esterhuysen, C.; Frenking, G. Chem. Eur. J. 2005, 11, 1813−1825. (50) (a) Cremer, D.; Kraka, E. Angew. Chem., Int. Ed. Engl. 1984, 23, 627−628. (b) Cremer, D.; Kraka, E. Croat. Chem. Acta 1985, 57, 1259−1281. (c) Kraka, E.; Cremer, D. Chemical Implication of Local Features of the Electron Density Distribution. In Theoretical Models of Chemical Bonding: The Concept of the Chemical Bond; Maksic, Z. B., Ed.; Springer-Verlag: New York, 1990. (51) (a) Glendening, E. D.; Badenhoop, J. K.; Reed, A. K.; Carpenter, J. E.; Bohmann, J. A.; Morales, C. M.; Weinhold, F. NBO 5.0; Theoretical Chemistry Institute, University of Wisconsin: Madison, WI, 2001. (b) Reed, A. E.; Curtiss, L. A.; Weinhold, F. Chem. Rev. 1988, 88, 899−926. (c) Weinhold, F.; Landis, C. R. Valency and Bonding: A Natural Bond Order Donor-Acceptor Perspective; Cambridge University Press: Cambridge, U.K., 2005. (52) Mercier, H. P. A.; Moran, M. D.; Schrobilgen, G. J.; Suontamo, R. J. J. Fluorine Chem. 2004, 125, 1563−1578. (53) Politzer, P.; Murray, J. S. General and Theoretical Aspects of the C−X Bonds (X = F, Cl, Br, I): Integration of Theory and Experiment. In The Chemistry of Functional Groups (Supplement D2); Patai, S., Rappoport, Z., Eds.; John Wiley & Sons: New York, 1995. (54) Landis, C. R.; Weinhold, F. J. Comput. Chem. 2007, 28, 198− 203. (55) Wu, Y.-D.; Peng, Z.-H.; Chan, K. W. K.; Liu, X.; Tuinman, A. A.; Xue, Z. Organometallics 1999, 18, 2081−2090. (56) (a) Huang, H.; Hurubeanu, N. R.; Bourgeois, C. J.; Cheah, S.-M.; Yuan, J.; Rheingold, A. L.; Hughes, R. P. Can. J. Chem. 2009, 87, 151− 160. (b) Leyssens, T.; Peeters, D. J. Org. Chem. 2008, 73, 2725−2730. (c) Leyssens, T.; Peeters, D.; Orpen, A. G.; Harvey, J. N. New J. Chem. 2005, 29, 1424−1430. (d) Leyssens, T.; Peeters, D.; Orpen, A. G.; Harvey, J. N. Organometallics 2007, 26, 2637−2645. (57) SMART 5.632; Bruker AXS, Inc.: Madison, WI, 2005. (58) SAINT 6.45; Bruker AXS, Inc.: Madison, WI, 2003. (59) Sheldrick, G. SADABS 2.05; University of Göttingen: Göttingen, Germany, 2002. (60) SHELXTL 5.10; Bruker AXS, Inc.: Madison, WI, 1997. (61) APEX II 1.08; Bruker AXS, Inc.: Madison, WI, 2004. (62) SAINT+ 7.06; Bruker AXS, Inc.: Madison, WI, 2003. (63) Jaguar 7.0−7.7; Schrödinger, LLC: New York, 2007−2010. (64) Keith, T. A. AIMAll, version 11.10.16; TK Gristmill Software: Overland Park, KS, 2011. [email protected].

(29) Brothers, P. J.; Burrell, A. K.; Clark, G. R.; Rickard, C. E. F.; Roper, W. R. J. Organomet. Chem. 1990, 394, 615−642. (30) King, R. B.; Bisnette, M. B. J. Organomet. Chem. 1964, 2, 15−37. (31) Taylor, R. C. J. Chem. Phys. 1954, 22, 714−716. (32) For representative examples of other titanium(IV) metallocene complexes containing a terminal Ti−F linkage, see: (a) [1,3-(SiMe3)2C5H3]2TiF2, Ti−F = 1.832(2) Å; [1,2,4-(SiMe3)3-C5H2][(SiMe3) C5H4]TiF2, Ti−F = 1.838(2) Å: Winter, C. H.; Zhou, X. X.; Heeg, M. J. Inorg. Chem. 1992, 31, 1808−1815. (b) [1,3-(SiMe3)2-C5H3][(SiMe3) C5H4]TiF2, Ti−F = 1.8357(9),1.855(2) Å: Winter, C. H.; Lewkebandara, T. S.; Shui, X.; Zhou, X. X. J. Crystallogr. Spectrosc. Res. 1993, 23, 685− 687. (c) (C5Me5)2Ti(CHCH2)F, Ti−F = 1.847(3) Å: Beckhaus, R.; Sang, J.; Oster, J.; Wagner, T. J. Organomet. Chem. 1994, 484, 179−190. (d) [(C5H4)2-C2Me4]TiF2, Ti−F = 1.855(1) Å: Edwards, A. J.; Burke, N. J.; Dobson, C. M.; Prout, K.; Heyes, S. J. J. Am. Chem. Soc. 1995, 117, 4637−4653. (e) [(C5H4)2-CMe2]TiF2, Ti−F = 1.8565(10), 1.8648(9)Å: Picka, M.; Cisarova, I.; Vinklarek, J.; Erben, M. Acta Crystallogr., Sect. E: Struct. Rep. Online 2005, E61, m1266−m1268. (33) For representative examples of other titanium(IV) complexes containing a single terminal Ti−OSO2CF3 linkage, see: (a) [(2-O-3,5Me2-C6H2CH2)3N]Ti(O3SCF3), Ti−OOTf = 2.002(2) Å: Bull, S. D.; Davidson, M. G.; Johnson, A. L.; Robinson, D. E. J. E.; Mahon, M. F. Chem. Commun. 2003, 1750−1751. (b) [ArNC(Me)] 2 CH] Ti(Me)2(O3SCF3), Ti−OOTf = 2.044(1) Å: Basuli, F.; Bailey, B. C.; Watson, L. A.; Tomaszewski, J.; Huffman, J. C.; Mindiola, D. J. Organometallics 2005, 24, 1886−1906. (c) [{ArNC(Me)}2CH]Ti( NSiMe3)(O3SCF3), Ti−OOTf = 2.0235(18) Å; [N{2-P(CHMe2)2-4-MeC6H3}2]Ti(NSiMe3)(O3SCF3), Ti−OOTf = 2.019(3) Å: Bailey, B. C.; Basuli, F.; Huffman, J. C.; Mindiola, D. J. Organometallics 2006, 25, 2725−2728. (d) [{ArNC(Me)}2CH]Ti(CHSiMe3)(THF)(O3SCF3), Ti−OOTf = 2.073(2) Å: Basuli, F.; Adhikari, D.; Huffman, J. C.; Mindiola, D. J. J. Organomet. Chem. 2007, 692, 3115−3120. (34) (a) (C5H4CH3)2Ti(CH2OCH3)(Cl): Erker, G.; Schlund, R.; K r u e g e r , C . J. Organomet. Chem. 1988, 338, C 4 −C6. (b) [(CH3)2Si(C5H4)2]Ti(CH3)(Cl): Gomez, R.; Cuenca, T.; Royo, P.; Hovestreydt, E. Organometallics 1991, 10, 2516−2518. (c) Cp2Ti[CH2Si(CH3)3](Cl): Eisch, J. J.; Caldwell, K. R.; Werner, S.; Krueger, C. Organometallics 1991, 10, 3417−3419. (35) Rodulfo de Gil, E.; De Burguera, M.; Valentina Rivera, A.; Maxfield, P. Acta Crystallogr., Sect. B 1977, B33, 578−579. (36) Single-crystal X-ray diffraction studies were performed on Cp2Ti(CF3)(Br) (13) and Cp2Ti(CF3)(I) (14), but accurate bond lengths could not be determined due to significant disorder between the CF3 and the Br/I ligands, respectively. (37) (a) Wilford, J. B.; Stone, F. G. A. Inorg. Chem. 1965, 4, 389− 393. (b) Clark, H. C.; Tsai, J.-H. J. Organomet. Chem. 1967, 2, 515− 517. (c) Cotton, F. A.; Wing, R. M. J. Organomet. Chem. 1967, 9, 511− 517. (38) Graham, W. A. G. Inorg. Chem. 1968, 7, 315−321. (39) Hall, M. B.; Fenske, R. F. Inorg. Chem. 1972, 11, 768−775. (40) (a) Bent, H. A. Chem. Rev. 1961, 61, 275−311. (b) Kaupp, M. Chem. Eur. J. 1999, 5, 3631−3643. (41) Loizou, D. C.; Castillo, J.; Oki, A. R.; Hosmane, N. S.; Morrison, J. A. Organometallics 1992, 11, 4189−4193. (42) Oberhammer, H. J. Fluorine Chem. 1983, 23, 147−162. (43) (a) Lee, C.; Yang, W.; Parr, R. G. Phys. Rev. B: Condens. Matter Mater. Phys. 1988, 37, 785−789. (b) Becke, A. D. J. Chem. Phys. 1993, 98, 5648−5652. (c) Becke, A. D. J. Chem. Phys. 1993, 98, 1372−1377. (d) Stephens, P. J.; Devlin, F. J.; Chabalowski, C. F.; Frisch, M. J. J. Phys. Chem. 1994, 98, 11623−11627. (e) Andersson, M. P.; Uvdal, P. J. Phys. Chem. A 2005, 109, 2937−2941. (44) (a) Dunning, T. H.; Hay, P. J. In Applications of Electronic Structure Theory; Schaefer, H. H., III, Ed.; Modern Theoretical Chemistry; Plenum: New York, 1977; Vol. 4. (b) Hay, P. J.; Wadt, W. R. J. Chem. Phys. 1985, 82, 299−310. (c) Hay, P. J.; Wadt, W. R. J. Chem. Phys. 1985, 82, 270−283. (d) Wadt, W. R.; Hay, P. J. J. Chem. Phys. 1985, 82, 284−298. (45) Numerous single-crystal X-ray diffraction studies were performed on Cp2Ti(CH3)(Cl), but accurate bond lengths could 1499

dx.doi.org/10.1021/om201055e | Organometallics 2012, 31, 1484−1499