Controlling the Reactivity of the Boronyl Group in Platinum Complexes

Oct 16, 2015 - Synopsis. A theoretical study of the cyclodimerization of (Cy3P)2Pt(BO)Br and [(Cy3P)2Pt(BO)]+ (Cy = cyclohexyl) suggests that the reac...
2 downloads 0 Views 1MB Size
Download PDF
Article pubs.acs.org/IC

Controlling the Reactivity of the Boronyl Group in Platinum Complexes toward Cyclodimerization: A Theoretical Survey Zhong Zhang,*,†,‡ Liang Pu,† Qian-shu Li,‡ and R. Bruce King*,§ †

College of Science, Northwest A&F University, Yangling, Shanxi 712100, People’s Republic of China Center for Computational Quantum Chemistry, South China Normal University, Guangzhou 510631, People’s Republic of China § Department of Chemistry and Center for Computational Chemistry, University of Georgia, Athens, Georgia 30602, United States ‡

S Supporting Information *

ABSTRACT: A theoretical study of the cyclodimerization of (Cy3P)2Pt(BO)Br (1Br) and [(Cy3P)2Pt(BO)]+ (1) (Cy = cyclohexyl) suggests that the reactivity of the BO ligand is primarily controlled by M←BO σ donation. Therefore, increasing the electron density at the metal center through strong σ-donor and weak π-acceptor ancillary ligands and a low formal metal oxidation state are suggested to reduce the polarity of the boronyl ligand and thus lower its reactivity toward cyclodimerization. The stable 1Br has lower Pt←BO σ donation and thus a less electrophilic boron atom, leading to a less polarized BO ligand. However, 1 is unstable in dichloromethane, since the dicationic dimer and transition state are highly stabilized by strong electrostatic interactions.

Co→BO π back-donation. This is consistent with an earlier theoretical study by Baerends and co-workers10 comparing the isolobal ligands BF, BNH2, BN(CH3)2, and BO− in their transition-metal complexes. Our recent theoretical study9 also predicts Co(BO)(CO)4 to be reactive toward cyclooligomerization to give complexes containing BnOn rings. However, other transition-metal systems have been shown to form stable boronyl complexes. Thus, the first metal boronyl complex (Cy3P)2Pt(BO)Br (Cy = cyclohexyl) was synthesized by Braunschweig and co-workers11 by the oxidative addition of Me3SiOBBr2 to (Cy3P)2Pt followed by Me3SiBr elimination. Subsequently, [Cp*Ru(μ2-H)]3(μ3-CPh)(μ3-BO) was obtained by Suzuki and co-workers by direct reaction of [Cp*Ru(μ2-H)]3(μ3-CPh)(μ3-BH) with H2O.12 A theoretical study of the synthesis of (Cy3P)2Pt(BO)Br indicated that the rate-limiting step is the Me3SiBr elimination. These observations suggested the mixed carbonyl boronyl complexes (PR3)2MClBr(BO) (CO) (M = Ir, Rh) as promising synthetic targets.13 However, the polarized BO ligand was found to act as a Lewis base or nucleophile at the electron-rich oxygen atom.14 In addition, the labile Br− anion in (Cy3P)2Pt(BO)Br is reactive toward nucleophilic substitution by PhS− and CH3CN.11,14 Interestingly, debromination of (Cy3P)2Pt(BO)Br with Ag+ led to cyclodimerization of the BO ligand in the initially generated [(Cy3P)2Pt(BO)]+ cation to give the [(Cy3P)2Pt2(B2O2)]2+ dication containing a B2O2 ligand bridging the two Pt atoms (Scheme 1).15 Such chemistry

1. INTRODUCTION Boron monoxide (BO) is a reactive radical forming diverse experimentally known boronyl-containing clusters in which the BO groups show remarkable structural and chemical integrity.1 Species containing BO units include the experimentally observed B(BO)2−, B(BO)3−, and the unusual BB-bonded B2(BO)22− as well as the theoretically predicted very stable B(BO)4−.2,3 The reactivity of free BO is related to the large electronegativity difference between boron and oxygen. This leads to a polarized π bond with the electrons largely localized on the oxygen atom. Because of this electronegativity difference, even the dimer B2O2 is a reactive species predicted to polymerize to graphene-like materials containing aromatic B3O3 rings.4 Moreover, high-energy multiply bonded BO and BN units are very reactive functional groups, leading to lowenergy σ-bonded products. For example, the organic BN double bond is reported to react reversibly with CO2.5 In this connection, pursuing stable organic and inorganic compounds containing boronyl groups is a potential source of new and useful chemistry. However, no stable organic boronyl compounds (RBO) have been reported, owing to the easy formation of trimers (R3B3O3) with aromatic boroxine rings.6,7 Transition-metal complexes are useful for stabilizing reactive small molecules through both electronic and steric effects, as exemplified by the well-known carbene complexes.8 The boronyl group in transition-metal complexes should be viewed as the anionic two-electron donor BO−, isoelectronic with cyanide, CN−. Our recent theoretical work9 indicated that the bonding of the anionic BO− ligand in Co(BO)(CO)4 involves very strong Co←BO σ donation combined with relatively weak © XXXX American Chemical Society

Received: July 20, 2015

A

DOI: 10.1021/acs.inorgchem.5b01597 Inorg. Chem. XXXX, XXX, XXX−XXX

Article

Inorganic Chemistry

work are divided into two different groups. The double-ζ plus polarization (DZP)26−29 basis sets were used for smaller [(PMe3)2Pt(BO)Br]n (n = 1, 2) complexes, except for the platinum atom, for which the effective core potential (ECP) basis set Lanl2DZ was used.30,31 However, the molecules containing PCy3 ligands are too large to use the DZP level basis set for all atoms. Therefore, the corresponding calculations for the larger complexes used mixed basis sets with minimum STO-3G basis sets32 for the outer-sphere Cy groups (i.e., C and H atoms). For the inner-sphere frameworks of these complexes (i.e., B, O, P, and Br) the DZP basis sets were used combined with the Lanl2DZ basis set for the platinum atoms. The geometries of all structures were fully optimized using the M06-L method and the basis sets mentioned above. Vibrational frequencies were determined by evaluating analytically the second derivatives of the energy with respect to the nuclear coordinates. All of the computations were carried out with the Gaussian 09 program,33 exercising the fine grid option (75 radial shells, 302 angular points) for evaluating integrals numerically.34 The tight designation (10−8 hartree) was the default for the self-consistent field (SCF) convergence. In order to consider solvent effects, the polarizable continuum models (PCM)35 were used to reoptimize structures (both local minima and the transition states) as well as the vibrational frequencies for predicting the free energy plots in dichloromethane solutions. Morokuma energy decomposition analyses (EDA)36 were carried out for selected BP86 structures at the BP86/DZP level using the ADF program package.37 The bonding dissociation energy is divided into two physically appealing entities:

Scheme 1. Debromination of Stable (Cy3P)2Pt(BO)Br Leading to Cyclodimerization

exemplifies controlling the reactivity of boronyl ligands in transition-metal complexes using electronic rather than steric effects.16 We now use theoretical methods to explore the reactivities of (Cy3P)2Pt(BO)Br and [(Cy3P)2Pt(BO)]+ toward cyclodimerization in connection with the syntheses of stable boronyl complexes and small B2O2 ring bridging complexes.

2. THEORETICAL METHODS Electron correlation effects were considered using density functional theory (DFT) methods, which have evolved as a practical and effective computational tool, especially for organometallic compounds.17−23 The functional used in this work is a hybrid meta-GGA DFT method, M06-L, developed by Truhlar’s group.24,25 The studies in Truhlar’s group suggest that M06-L is one of the best functionals for the study of organometallic and inorganic thermochemistry and is perhaps the best currently available functional for transition-metal energetics. DFT methods are fortunately less basis set sensitive than higherlevel methods such as coupled cluster theory. Basis sets used in this

Figure 1. Optimized structures in both the gas phase (values in black) and dichloromethane solvent (values in red). Bond distances are given in Å. B

DOI: 10.1021/acs.inorgchem.5b01597 Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry

Article

ΔE = ΔEprep + ΔE int

Table 1. Energy Decomposition Analysis of MeBr1, Me1, and Co(CO)4(BO) (in kcal/mol)

The preparation energy (ΔEprep) is the energy required to promote the two fragments (A and B) from their isolated equilibrium geometries to their geometries in the compound AB. The interaction energy (ΔEint) is the interaction between the two prepared fragments in the molecule. In this study, ΔEprep is sufficiently small to be neglected, so that only ΔEint needs to be considered. This interaction energy can be separated into three major components:

ΔEint ΔEPauli ΔEels ΔEster ΔEorb a1a b1a b2a

ΔE int = ΔEels + ΔE Pauli + ΔEorb In this equation the first term, ΔEels, is the electrostatic interaction energy between the fragments, which is calculated using a frozen electron density distribution at the geometry of the complex. The second term, ΔEPauli, is the repulsive energy caused by Pauli repulsion. The last term, ΔEorb, is the stabilization energy from relaxed orbital interaction between the prepared fragments. The ΔEorb term can be broken down further into orbital contributions from different irreducible representations. This allows the prediction of separate energy contributions from σ and π interactions.

MeBr1

Me1

Co(CO)4(BO)

−214.8 318.8 −402.7 −83.9 −126.6 −104.9 −10.3 −10.4

−343.3 371.2 −543.6 −172.4 −166.6 −138.6 −12.7 −13.6

−252.2 260.9 −335.3 −74.4 −175.1 −152.5 −22.4(e1) −0.2

a The a1 term represents the M←L σ donation, and the b1 + b2 (e1) term represents the M→L π* back-donation.

(23.3 ± 3.5 kcal/mol). This indicates the primary importance of the σ donation. Weaker M←BO σ donations provide less electron density to the metal center, thereby leading to a less electrophilic boron atom and thus a less polarized boronyl, which is less reactive toward cyclodimerization. However, the σ donation can be effectively reduced by increasing the electron density at the metal center using ancillary ligands that are strong σ donors as well as weak π acceptors. Therefore, removing the soft Br− ligand from neutral 1Br (with one Br− and two PCy3 ligands), leading to cationic 1, reduces slightly the natural platinum negative charge from −0.61 in 1Br to −0.59 in 1 for weakening the M←BO σ donation. Consequently, the Pt←BO σ donation in 1Br is decreased to −104.9 kcal/mol, thus providing a BO bond less polarized (natural charges of 0.65 for B, − 0.98 for O, and thus −0.33 for BO) than the more polarized BO bond in 1 (natural charges of 0.75 for B, − 0.87 for O, and thus −0.12 for BO). Alternatively, the Pt−B(O) interaction can be viewed as a covalent bond between a BO radical and a platinum fragment. NBO analysis41 indicates that such a localized covalent Pt−B bond in 1Br originates from a relatively smaller contribution of 42.4% from the platinum atom and a larger contribution of 57.6% from the boron atom. However, the covalent electron pair of the Pt−B bond in cationic 1 is shifted to the platinum side, as indicated by a larger contribution of 57.7% from the platinum atom and a smaller 42.3% from the boron atom. Such a covalent electron pair redistribution arises from the cationic 1 having a more electrophilic platinum atom than 1Br. Thus, stable 1Br has a less positively charged boron atom as well as more total negative charge to prevent dimerization in comparison with the cationic I. However, we prefer to consider metal boronyl complexes as containing the anionic BO− ligand isoelectronic with the neutral CO ligand in metal carbonyls. This interpretation is supported by the negative charges of −0.33 and −0.12 on the BO group in 1Br and 1, respectively. Relative to 1 or 1Br with strong PCy3 Lewis bases, the strong π acid carbonyl ligands42 in Co(CO)4(BO) increase the M←BO σ donation, resulting in a more polarized BO bond (natural charges of 1.09 for B, − 0.83 for O, and thus 0.26 for BO). In this connection, the experimentally known neutral (Cy3P)2PtBr(BO) is stabilized by Cy3P and soft Br− ligands but the neutral Co(CO)4(BO) is destabilized by the strong Lewis π-acceptor carbonyl ligands. Furthermore, N-heterocyclic carbenes (NHCs), which are stronger σ donors than PR3 and weak π acceptors, should be reasonable ligand(s) for stabilizing boronyl metal complexes toward dimerization.43 Moreover, the BO bond lengths in 1Br, 1, and Co(CO)4(BO) are predicted to be 1.233, 1.226, and 1.218 Å, respectively, in the gas phase.

3. RESULTS AND DISCUSSION The experimentally known (Cy 3 P) 2 Pt(BO)Br and [(Cy3P)4Pt2(B2O2)]2+ have been characterized by X-ray diffraction. However, the geometric parameters for (Cy3P)2Pt(BO)Br determined by X-ray diffraction are unreliable owing to disorder. We therefore compare our predicted [(Cy3P)4Pt2(B2O2)]2+ geometric parameters with the X-ray diffraction results (Figure 1 and Table S1 in the Supporting Information). The average experimental Pt−B and B−O lengths are found to be 1.950 and 1.413 Å,15 respectively, which are very close to the gas-phase results (1.952 and 1.416 Å) as well as the results in dichloromethane (1.946 and 1.417 Å). Both the optimized gas-phase and dichloromethane solution structures are very similar. Therefore, only the gasphase results are mentioned, except for a few significant differences. Agostic C−H−Pt bonds involving cyclohexyl hydrogens are found in the optimized [(Cy3P)2Pt(BO)]+ and [(Cy3P)4Pt2(B2O2)]2+ structures, as indicated by relatively short Pt−H distances (Figure 1). However, both experimental and theoretical Pt−H bond distances have limited accuracy, since X-ray diffraction cannot locate precisely the hydrogen atoms and the STO-3G basis used for the ligand carbon and hydrogen atoms cannot accurately predict the agostic bonds. However, such inaccuracies are expected to have a minor effect on our prediction of boronyl behavior because of the weakness of the agostic C−H−Pt bonds in 1 and 2.38 Each monomer, i.e. 1Br and 1 (Figure 1), has a triply bonded boronyl group, as indicated by molecular orbitals.11 In order to reveal specific bonding contributions, energy decomposition analysis (EDA) 36 has been applied to the simplified (Me 3 P) 2 PtBr(BO), and [(Me 3 P) 2 Pt(BO)] + structures (1MeBr and 1Me in Figure 1). The large interaction energies (Eints) for these boronyl complexes, which are more than −214.8 kcal/mol, indicate that the charge-transfer interaction contributes to strong electrostatic interaction for M←BO bonding. Thus, the attractive electrostatic interactions (Eelss) are more than −330 kcal/mol between cationic [M]+ fragments and anionic BO−. However, the orbital interactions (Eorbs), of more than 126 kcal/mol (about half of Eelss), might also play an important role in the M←BO interactions (Table 1). On the basis of the popular Dewar−Chatt−Duncanson (DCD) model,39,40 the Eorbs originate from large contributions of M←BO σ donation energies (from −104.9 to −152.5 kcal/ mol), as well as those of smaller M→BO π back-donation(s) C

DOI: 10.1021/acs.inorgchem.5b01597 Inorg. Chem. XXXX, XXX, XXX−XXX

Article

Inorganic Chemistry Apparently, the M→BO π back-donation decreases in the sequence 1Br > 1 > Co(CO)4(BO). However, the EDA results (Table 1) for Co(CO)4(BO), 1MeBr, and 1Me show that the M→BO π back-donation decreases in the sequence 1Me > Co(CO)4(BO) > 1MeBr, even though 1MeBr has the longest BO distance of 1.228 Å (Figure 1), consistent with a large BO Wiberg bond index of 1.86 in 1MeBr relative to a smaller BO Wiberg bond index of 1.74 in 1Br. This contradiction may arise from the decreased ionic character of the BO bond, as well as increased repulsion between the negatively charged Pt and O atoms in (Me3P)2PtBr(BO). In short, for stabilizing the boronyl group in transition-metal complexes, we therefore suggest that ancillary ligands that are strong σ donors and weak π acceptors are effective for reducing the polarity of the boronyl ligand. The earlier reported possible synthetic targets, (Me3P)2MClBr(BO) (CO) (M = Ir, Rh),13 contain a strong π-acceptor CO ligand42 and an M(III) center, which increase the polarity of the BO group as a consequence of the increased M←BO σ donation. We predict at the M06L/DZP(Lanl2dz) level that monomeric (Me3P)2IrClBr(BO)(CO) (1Ir in Figure 1) is disfavored with respect to formation of the corresponding dimer (2Ir in Figure 1). On comparison of the highest ν(CO) frequency (2094 cm−1) for Fe2(CO)6 at the B3LYP/DZP level,44 the even higher ν(CO) frequency (2126 cm−1) in 1Ir suggests weaker Ir→CO π back-donation relating to the high Ir(III) formal oxidation state. Therefore, the Ir−Br distance in 1Ir of 2.617 Å (Figure 1) is 0.057 Å shorter than the Pt−Br distance in 1Br. The obviously reduced Ir−Br distance mainly arises from strong IrIII←Br− donation, leading to less negatively charged Br− ligands (−0.27) in 1Ir than in 1Br (−0.58). Analogously, the IrIII←BO− σ donation is also enhanced, resulting to a more polarized BO bond (natural charges of 1.02 for B, − 0.89 for O, and thus 0.13 for BO). As a consequence, the more polarized boronyl in 1Ir is predicted to be disfavored toward cyclodimerization into 2Ir by releasing a Gibbs free energy of 15.8 kcal/mol. These observations illustrate how metal centers in low formal oxidation states are effective for reducing the polarity of the boronyl ligand and thus the reactivity of the boronyl ligand toward cyclodimerization. The predicted structures of the cyclodimers, namely 2Br and 2, are shown in Figure 1. Each dimer is found to have a B2O2 ring forming a B→Pt bond with each platinum atom. For the B2O2 ring in 2, our predicted average B−O bond distance of 1.416 ± 0.001 Å perfectly matches the experimental value for the B2O2 ring in [(Cy3P)2Pt2(B2O2)]2+ with an average B−O distance of 1.413 Å.15 The B−O distances of the B2O2 ring in 2Br increase to 1.425 ± 0.013 Å, possibly owing to electron donation from Br− to Pt to reduce the positive charge on boron, thereby reducing the natural charge of 0.84 on the boron in 2Br relative to the 0.92 for the boron in 2 (Table S2 in the Supporting Information). This reduces the ionic character of the B−O bond and thus its strength. As a result, the four B−O bonds in the B2O2 ring should be B−O single bonds, as suggested by their Wiberg bond indices of 0.70 ± 0.01. The relatively short B···B distance of 1.847 ± 0.02 Å in the dimers originates from the four-membered B2O2 ring structure and repulsion of the lone pairs on the two oxygen atoms, analogous to the Si2O2 ring.45 The bromide dimer 2Br is reported to have two Pt−B distances of 1.986 ± 0.001 Å, which are also elongated ∼0.035 Å from the distances in 2, predicted in the gas phase. Note that the agostic C−H−Pt bonds in 2 and Pt−Br bonds in 2Br are weakened in dichloromethane. In

particular, one of the Pt−Br bond distances increases from the original value of 2.793 Å in 2Br in the gas phase to a nonbonding value of 5.159 Å in the solvent, indicating the lability of the Br− ligand. Solvent effects do not significantly change the predicted structures for both the monomers and dimers. However, the results in Figure 2 show that the polar dichloromethane solvent

Figure 2. Cyclodimerization mechanisms. The upper (black) route refers to the gas-phase reactions, whereas the lower (red) route represents reactions in the dichloromethane solvent. The relative Gibbs free energies are given in kcal/mol.

greatly affects the stability of the boronyl complexes with respect to cyclodimerization. Formation of 2Br from two separate 1Br units is thermodynamically disfavored owing to the positive predicted Gibbs free energies (ΔG) of 26.0 kcal/ mol for the reaction in the gas phase and 17.3 kcal/mol in dichloromethane. In the gas phase, formation of 2 from two 1 units is a possible fluxional process, since the predicted Gibbs free energy for this dimerization is only −0.8 kcal/mol. However, the dimer 2 is highly favored in dichloromethane because of the strongly negative −18.6 kcal/mol Gibbs free energy of formation. This means that the dication 2 is highly stabilized by the polar solvent, possibly owing to strong electrostatic interaction between solute and solvent. Indeed, the solvation energies for 2 and 2Br are predicted to be −40.3 and 16.7 kcal/mol, respectively, primarily originating from the electrostatic terms (Ges) of −82.3 kcal/mol for 2 and −21.0 kcal/mol for 2Br. These dimers originate from dimerization of two 1Br or 1 complexes through the positively charged boron atom and negatively charged oxygen atom in one BO bond of one monomer attracting the oxygen and boron atoms, respectively, in the BO bond of the other monomer. Related transition states (2Brts and 2ts in Figure 1) have been located for determining the energy barrier of this dimerization. By following the corresponding imaginary vibrational frequencies, forward IRC calculations lead to the cyclodimer products 2 and 2Br. Each transition state is found to have a B2O2 trapezoid with a short B···O distance of 1.633 Å for 2Brts and 1.828 Å for 2ts and a long B···O distance of 2.555 Å for 2Brts and 2.679 Å for 2ts. The relatively short B···O distance originates from the LUMO of one 1 or 1Br (mainly located on the boron atom) accepting HOMO electrons (lone pair electrons located on the oxygen atom) from the other 1 or 1Br (see Figure S1 in the Supporting Information). This interaction converts the boron hybridization from the original linear sp in the monomer to approximately sp2 with an ∠OBO angle of 110.4 ± 1.0° in 2ts or 2Brts. In addition, the directly bonded B−O distances are predicted to be 1.254 ± 0.002 Å for 2ts and 1.277 ± 0.002 Å for 2Brts, which are ∼0.028 and ∼0.044 Å longer than the triply bonded BO distances in 1 and 1Br, respectively. Such D

DOI: 10.1021/acs.inorgchem.5b01597 Inorg. Chem. XXXX, XXX, XXX−XXX

Inorganic Chemistry

■

elongations of BO distances indicate weakening of the B−O bond, as also suggested by reduced Wiberg bond indices from 1.74 in 1 to 1.42 in 2 and from 1.59 in 1Br to 1.32 in 2Br. Dimerization of 1Br is kinetically unfavorable, owing to the very high 56.0 kcal/mol Gibbs free energy of activation in the gas phase and a somewhat lower value of 48.8 kcal/mol in dichloromethane (Figure 2). 2Brts is slightly stabilized in dichloromethane because it is predicted to have a larger dipole moment of 29.9 D, providing a dipole−dipole interaction stronger than that of 11.8 D in the gas phase. In contrast, dimerization of 1 is more feasible, since the corresponding Gibbs free energies of activation are found to be 34.0 kcal/mol in the gas phase and only 14.2 kcal/mol in dichloromethane (Figure 2). 2ts is therefore more stabilized than 1 in dichloromethane, as suggested by its solvation energy of −82.2 kcal/mol for the 2ts dication and −33.4 kcal/mol for 1 (Table S3 in the Supporting Information). In addition, the dimerization processes are highly controlled by entropy, since the Gibbs free energies of reaction are generally larger than the ZPE corrected total electronic energies by more than 20.0 kcal/ mol. For instance, the ZPE energies and Gibbs free energies of reaction for formation of 2 are found to be 5.4 (ΔEZPE) or 26.0 kcal/mol (ΔG) for 2ts and 36.0 (ΔEZPE) or 56.0 kcal/mol (ΔG) for 2 in the gas phase, respectively. However, the threecoordinate intermediate 1 can be stabilized as the fourcoordinate [(Cy3P)2Pt(BO)(NCCH3)]+ in acetonitrile solution,14 even though 1 is highly susceptible to cyclodimerization in other polar solvents. In summary, the cyclodimerization of neutral 1Br to 2Br is thermodynamically and kinetically unfavorable in both the gas phase and dichloromethane solution. However, the dimerization of cationic 1 to dicationic 2 is feasible owing to stabilization of the 2ts dication and 2.

AUTHOR INFORMATION

Corresponding Authors

*E-mail for Z.Z.: [email protected]. *E-mail for R.B.K.: [email protected]. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS We are indebted to the National Natural Science Foundation of China (Grant 21303138), the China Postdoctoral Science Foundation (Grants 2013M540660 and 2014T70817), the Chinese Universities Scientific Fund (Grant 2014YB028), and the U.S. National Science Foundation (Grant CHE-1057466) for support of this research.

■

REFERENCES

(1) Zhai, H.-J.; Chen, Q.; Bai, H.; Li, S.-D.; Wang, L.-S. Acc. Chem. Res. 2014, 47, 2435−2445 and references therein. (2) Zhai, H.-J.; Li, S.-D.; Wang, L.-S. J. Am. Chem. Soc. 2007, 129, 9254−9255. (3) Li, S.-D.; Zhai, H.-J.; Wang, L.-S. J. Am. Chem. Soc. 2008, 130, 2573−2579. (4) Zhang, Z.; Pu, L.; Li, Q.-S.; King, R. B. Inorg. Chem. 2015, 54, 2910−2915. (5) Xie, L.; Zhang, J.; Hu, H.; Cui, C. Organometallics 2013, 32, 6875−6878. (6) Kinney, C. R.; Pontz, D. F. J. Am. Chem. Soc. 1936, 58, 197−197. (7) Parker, D. S. N.; Dangi, B. B.; Balucani, N.; Stranges, D.; Mebel, A. M.; Kaiser, R. I. J. Org. Chem. 2013, 78, 11896−11900. (8) Crabtree, R. H. The Organometallic Chemistry of The Transition Metals, 4th ed.; Wiley: Hoboken, NJ, 2005. (9) Zhang, Z.; Pu, L.; Li, Q.-S.; King, R. B. RSC Adv. 2015, 5, 35558− 35563. (10) Ehlers, A. W.; Baerends, E. J.; Bickelhaupt, F. M.; Radius, U. Chem. - Eur. J. 1998, 4, 210−221. (11) Braunschweig, H.; Radacki, K.; Schneider, A. Science 2010, 328, 345−347. (12) Kaneko, T.; Takao, T.; Suzuki, H. Angew. Chem., Int. Ed. 2013, 52, 11884−11887. (13) Zeng, G.; Sakaki, S. Inorg. Chem. 2012, 51, 4597−4605. (14) Braunschweig, H.; Radacki, K.; Schneider, A. Chem. Commun. 2010, 46, 6473−6475. (15) Braunschweig, H.; Radacki, K.; Schneider, A. Angew. Chem., Int. Ed. 2010, 49, 5993−5996. (16) Brand, J.; Braunschweig, H.; Sen, S. S. Acc. Chem. Res. 2014, 47, 180−191. (17) Brynda, M.; Gagliardi, L.; Widmark, P. O.; Power, P. P.; Roos, B. O. Angew. Chem., Int. Ed. 2006, 45, 3804−3807. (18) Sieffert, N.; Bühl, M. J. Am. Chem. Soc. 2010, 132, 8056−8070. (19) Schyman, P.; Lai, W.; Chen, H.; Wang, Y.; Shaik, S. J. Am. Chem. Soc. 2011, 133, 7977−7984. (20) Adams, R. D.; Pearl, W. C.; Wong, Y. O.; Zhang, Q.; Hall, M. B.; Walensky, J. R. J. Am. Chem. Soc. 2011, 133, 12994−12997. (21) Lonsdale, R.; Olah, J.; Mulholland, A. J.; Harvey, J. N. J. Am. Chem. Soc. 2011, 133, 15464−15474. (22) Crawford, L.; Cole-Hamilton, D. J.; Drent, E.; Bühl, M. Chem. Eur. J. 2014, 20, 13923−13926. (23) Zhekova, H.; Krykunov, M.; Autschbach, J.; Ziegler, T. J. Chem. Theory Comput. 2014, 10, 3299−3307. (24) Zhao, Y.; Truhlar, D. G. J. Chem. Phys. 2006, 125, 194101− 194118. (25) Zhao, Y.; Truhlar, D. G. Theor. Chem. Acc. 2008, 120, 215−241. (26) Dunning, T. H. J. Chem. Phys. 1970, 53, 2823−2833. (27) Dunning, T. H.; Hay, P. J. Methods of Electronic Structure Theory; Schaefer, H. F., III, Ed.; Plenum: New York, 1977. (28) Huzinaga, S. J. Chem. Phys. 1965, 42, 1293−1302. (29) Wachters, A. J. H. J. Chem. Phys. 1970, 52, 1033−1036.

4. CONCLUSIONS Boronyl groups in transition-metal complexes are stabilized by M←BO σ donation to reduce the polarity of the B−O bond. Therefore, increasing the electron density at the metal center through strong σ-donor and weak π-acceptor ancillary ligands and a low formal metal oxidation state are suggested to reduce the polarity of the boronyl ligand and thus lower its reactivity toward cyclodimerization. The neutral 1Br is thermodynamically and kinetically stable toward cyclodimerization in both the gas phase and dichloromethane solution. In contrast, the cation 1 is disfavored with respect to dimerization to 2 in dichloromethane, since the 2ts and 2 dications are highly stabilized by strong electrostatic interactions between the dications and the polar solvent.

■

Article

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.inorgchem.5b01597. Comparison of selected experimental bond distances with theoretical results, some predicted natural charges for monomers, dimers, and transition states, specific solvation energy contributions for monomers, dimers, and transition states, the theoretical Cartesian coordinates for monomers (seven structures), dimers (five structures), and transition states (four structures) in the gas phase and in dichloromethane, and the complete ref 33 (PDF) E

DOI: 10.1021/acs.inorgchem.5b01597 Inorg. Chem. XXXX, XXX, XXX−XXX

Article

Inorganic Chemistry (30) Wadt, W. R.; Hay, P. J. J. Chem. Phys. 1985, 82, 284−298. (31) Hay, P. J.; Wadt, W. R. J. Chem. Phys. 1985, 82, 299−310. (32) Collins, J. B.; Schleyer, P. v. R.; Binkley, J. S.; Pople, J. A. J. Chem. Phys. 1976, 64, 5142−5151. (33) Frisch, M. J., et al. Gaussian 09, Revision B.01, Gaussian, Inc., Wallingford, CT, 2010. (34) Papas, B. N.; Schaefer, H. F., III J. Mol. Struct.: THEOCHEM 2006, 768, 175−181. (35) Foresman, J. B.; Keith, T. A.; Wiberg, K. B.; Snoonian, J.; Frisch, M. J. J. Phys. Chem. 1996, 100, 16098−16104. (36) te Velde, G.; Bickelhaupt, F. M.; Baerends, E. J.; Fonseca Guerra, C.; van Gisbergen, S. J. A.; Snijders, J. G.; Ziegler, T. J. Comput. Chem. 2001, 22, 931−967. (37) Fonseca Guerra, C.; Snijders, J. G.; te Velde, G.; Baerends, E. J. Theor. Chem. Acc. 1998, 99, 391−403. (38) Niu, S.; Hall, M. B. Chem. Rev. 2000, 100, 353−405. (39) Dewar, M. J. S. Bull. Soc. Chim. Fr. 1951, 18, C79−C83. (40) Chatt, J.; Duncanson, L. A. J. Chem. Soc. 1953, 2939−2947. (41) Carpenter, J. E.; Weinhold, F. J. Mol. Struct.: THEOCHEM 1988, 169, 41−62. (42) Diefenbach, A.; Bickelhaupt, F. M.; Frenking, G. J. Am. Chem. Soc. 2000, 122, 6449−6458. (43) Boehme, C.; Frenking, G. Organometallics 1998, 17, 5801− 5809. (44) Xie, Y.; Schaefer, H. F., III; King, R. B. J. Am. Chem. Soc. 2000, 122, 8746−8761. (45) Grev, R. S.; Schaefer, H. F., III J. Am. Chem. Soc. 1987, 109, 6577−6585.

F

DOI: 10.1021/acs.inorgchem.5b01597 Inorg. Chem. XXXX, XXX, XXX−XXX