Nature of Intramolecular Metal–Metal Interactions in Supported Group

Article pubs.acs.org/Organometallics

Nature of Intramolecular Metal−Metal Interactions in Supported Group 4−Group 9 and Group 6−Group 9 Heterobimetallic Complexes: A Combined Density Functional Theory and Topological Study Ujjal Gogoi, Ankur Kanti Guha, and Ashwini K. Phukan* Department of Chemical Sciences, Tezpur University, Napaam 784028, Assam, India S Supporting Information *

ABSTRACT: Quantum chemical calculations have been carried out on a series of supported group 4−group 9 and group 6−group 9 heterobimetallic complexes designated by the general formulas [Cp 2 M(μ-E) 2 M′(H)(CO)L] and [(CO)4M(μ-E)2M′(H)(CO)L] where E = SH, SeH or PH2 and L = PH3, CO, NHC, or aNHC. An analysis of the optimized geometries of these molecules indicates the presence of an M···M′ interaction. The nature of this interaction is investigated by using Bader’s quantum theory of atoms in molecules (QTAIM), electron localization function (ELF), and source function (SF). The results of QTAIM analysis suggest a polar covalent interaction between the two disparate metal centers in these heterobimetallic complexes. ELF analysis identifies a bonding basin between the two metal centers, while SF analysis reveals that the metal−metal bonding is moderately delocalized. The strength of the M···M′ interaction is found to be stronger in group 4−group 9 heterobimetallic complexes compared to group 6−group 9 ones. late heterobimetallic complexes and proposed the presence of a weak and dative metal−metal interaction. 8 In another work, they also proposed the presence of a dative interaction between the two metal centers in a thiolato-bridged heterobimetallic complex of V and Cu.9 In an extensive theoretical study, Jansen et al. have studied the metal−metal bond polarity in unsupported Ti−Co and Zr−Co heterobimetallic complexes.10 Their results reveal a highly polar character of the metal−metal bond, which is also effected by the nature of the trans-axial ligand coordinated to the late metal center, i.e., cobalt. Ishii et al. synthesized a series of heterodinuclear nitrosyl complexes and explained their reactivity on the basis of a dative metal− metal interaction.11 In recent times, Thomas et al. have carried out a series of work on heterobimetallic complexes featuring a metal−metal multiple bond between Zr and Co.12 They also studied the effect of substituents on metal−metal interactions in Zr−Pt heterobimetallics.13 Frenking et al. have carried out a systematic study on bonding in unbridged early late heterobimetallic complexes of group 4 and group 9 metals14 and metal−metal multiple-bonded systems of homo- and heterobimetallic complexes.15 In their first work, they studied the nature of metal−metal interactions and the influence of rotation of the terminal ligands on it, while in the other work,

1. INTRODUCTION In the realm of coordination chemistry, heterobimetallic complexes occupy a special position due to the presence of two different metal centers. The combination of two different metals introduces asymmetry and polarity within the complex as well as variation in electronic environment at the two metal centers, which enhances the reactivity of these complexes in comparison to their monometallic and homobimetallic counterparts.1−3 Such enhanced reactivity renders these complexes suitable for their use as catalysts in a variety of chemical reactions.1,2 Hydrogenation of alkenes4 and ketones,5 hydroformylation,3a polymerization,6 isomerization,4 and hydrosilylation7 of alkenes are some of the reactions in which heterobimetallic complexes are used as catalysts. It is also observed that the removal or replacement of the early metal fragment in a given heterobimetallic complex by a nonmetallic fragment reduces the efficiency or selectivity of the catalyst.1 This indicates that the presence of one metal atom modifies the properties of the other and, hence, changes the reactivity as well. Thus, the reactivity of heterobimetallic complexes might be influenced by the nature and extent of metal−metal interaction. A large number of studies have been devoted to reveal the nature and extent of metal−metal interactions in different heterobimetallic complexes.8−16 Stephan et al. have carried out extended Hückel and Fenske−Hall MO calculations on thiolato-bridged early © 2011 American Chemical Society

Received: September 5, 2011 Published: October 21, 2011 5991

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they studied the bonding situations in homo- and heterobimetallic complexes of group 6 elements. Thus, a variety of unsupported heterobimetallic complexes have been studied by different groups, and in most of them, the presence of a metal− metal interaction is either indicated or confirmed. However, to the best of our knowledge, to date, there exists no systematic study on the nature of metal−metal interactions in supported heterobimetallic complexes as a function of changes in the nature of the bridging and ancillary ligands. We present here a systematic theoretical study on the nature and extent of metal−metal interactions on a series of group 4− group 9 and group 6−group 9 heterobimetallic complexes (Scheme 1). The model complexes are based on the

wanted to check how replacement of PH3 by better σ-donating ligands will effect the extent of M···M′ interaction.

2. COMPUTATIONAL DETAILS All the molecules are optimized with gradient-corrected hybrid density functional theory (DFT) using Becke’s three-parameter hybrid functional with correlation from Lee, Yang, and Parr (B3LYP) at a temperature of 298.15 K and a pressure of 1.0 atm under gaseous conditions.18 We have used the LANL2DZ basis set with the effective core potential (ECP) of Hay and Wadt.19 Frequency calculations were performed at the same level of theory to characterize the nature of the stationary point. The ground states of all the structures were confirmed by verifying their respective Hessian (matrix of analytically determined second-order energy derivative) to be all real. Natural bonding analysis (NBO) was also performed to carry out the bonding analysis of the structures.20 All these calculations were performed using the Gaussian 03 suite of program.21 The nature of bonding in these complexes was analyzed with the help of Bader’s quantum theory of atoms in molecules (QTAIM).22 For QTAIM analysis, we have first generated the wave function from the optimized structures using Gaussian 03 and then analyzed with the AIMALL package.23 For generation of the wave functions, we have used the all-electron basis set DGDZVP (density gauss double-ζ with polarization function).24 Electron localization function calculations were performed with TopMod09,25 and source functions were calculated using a code provided by Carlo Gatti.26

Scheme 1

3. RESULTS AND DISCUSSION 3.1. Molecular Geometry. 3.1.1. Group 4−Group 9 Heterobimetallic Complexes. The optimized geometrical parameters for Ti−Co heterobimetallics are given in Table 1 (for geometrical parameters of Zr−Rh and Hf−Ir heterobimetallics, see Table S1 and Table S2, respectively, in the Supporting Information). The minimum energy structures of these molecules show a four-membered ring defined by the two metals and the two bridging ligands that is nearly planar in the case of SH- and SeH-bridged complexes, while it takes a butterfly shape in the case of PH2-bridged complexes (Figure 1). The butterfly shapes of the phosphido-bridged complexes are consistent with the experimental geometry3 with E−M′−E−M dihedral angles varying from 2.8° to 5.4°. The computed geometrical parameters of the model system [Cp2Zr(μ-PH2)2Rh(H)(CO)PH3] are in excellent agreement with the reported experimental structure [Cp2Zr(μ-PPh2)2Rh(H)(CO)PPh3].3a The calculated Zr−Rh distance (2.992 Å) is very close to that observed experimentally (2.980 Å). Also, the calculated angles ∠Zr−P− Rh (71.5°), ∠P−Zr−P (100.3°), ∠P−Rh−P (116.4°), and ∠Zr−Rh−Caxial (102.8°) are in reasonable agreement with the experimental values (72.4°, 95.6°, 118.3°, and 96.7° respectively). However, the calculated Rh−P distances (2.425 Å) differ from the experimental ones (2.320 Å). This might be due to the presence of two phenyl groups attached to the two bridging phosphorus atoms in the experimental structure, which we have modeled with hydrogen atoms. Moreover, the calculated Zr−Rh distances are very close to other experimentally observed supported Zr−Rh bimetallic complexes (range of Zr−Rh distance is 2.863−2.980 Å).3b In the case of SH- and SeH-bridged complexes, the substituents at the bridging atoms may adopt either a cisoid-endo conformation or a transoid one, as shown in Scheme 2.8 However, we have found very small energy differences ( W, which is in accord with the relative electronegativity of these elements. However, the charges at the late metal center are more or less similar to those obtained for group 4−group 9 systems. Thus, the M···M′ interactions of group 6−group 9 heterobimetallics are more polar than the group 4−group 9 ones. There is a slight increase in negative charge at the early metal center as the ancillary ligands are changed from PH3 through aNHC. However, this change in charge is appreciable at the late metal center. The bridging groups are found to have a profound effect on the charges at both the metal centers. Replacement of the chalcogenide groups by the phosphido group at the bridging position significantly increases the negative charge at both M and M′. In order to investigate the bonding situations in these heterobimetallic complexes, we have chosen a Ti−Co complex with E = SH and L = PH3 (1a) and a Cr−Co complex with E = SH and L = PH3 (1a′) as the representative molecules for group 4−group 9 and group 6−group 9 heterobimetallics, respectively. The HOMO (−5.1 eV) of 1a (Figure 5a) represents the σ-bonding interaction between the Ti and Co atoms with Co contributing more than Ti. The higher contribution of the late metal in the HOMO explains its lower positive charge. The LUMO (−2.0 eV) of 1a corresponds to the σ-antibonding 5996

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interaction between the two metal atoms with more contribution coming in from Ti. The HOMO (−6.5 eV) of 1a′ (Figure 5b) represents the σ-bonding interaction between Cr and Co atoms where the contribution of Cr is more. This higher contribution is in tune with high negative charges at the group 6 metals. It is evident from the shape of the HOMO of 1a and 1a′ that the extent of overlap between the two metal orbitals is larger in the former. This might be due to the presence of two Cp rings attached to the group 4 metals (Figure 5a), which forces the d orbital on Ti to lie along the Ti−Co axis. Such steric congestion was absent in group 6 fragments (Figure 5b), and as a result, the d orbital on group 6 metals tilts away from the Cr−Co axis. This results in a weaker overlap, and hence, the M···M′ interactions are weaker in group 6−group 9 heterobimetallics than in group 4− group 9 ones. Like 1a, the LUMO (−3.7 eV) of 1a′ is σ antibonding with respect to Cr and Co. Since, the LUMO is much lower lying, reduction of 1a′ may result in rupture of the M···M′ bond. Thus, the experimental observation of the rupture of the W−Ir bond of [(CO)4W(μ-PPh2)2Ir(H)(CO)PPh3] upon reduction to yield the dihydride complex [(CO)4W(μPPh2)2Ir(H)2(CO)PPh3] may be traced to the presence of a lower lying σ-symmetric W−Ir antibonding orbital.16c 3.3. QTAIM Analysis. Applying Bader’s quantum theory of atoms in molecules (QTAIM), it is possible to obtain a lot of important information about different chemical properties in a molecule.22 We have analyzed the topology of electron density using this theory. This theory is mainly based on the threedimensional electron density function, ρ(r). The topological analysis is the investigation of critical points of this function, ρ(r). Critical points are of three types, which vary with their ranks (number of nonzero eigenvalues λ i) and signature (sum of the sign of the eigenvalues). Among them, the most relevant one to our study is the bond critical point (BCP). A bond critical point is the point present along the bond path between two atoms with minimum electron density. It helps in predicting the presence of a bonded interaction between two atoms.22 The parameters that are commonly used to ascertain the nature and extent of bonding between two atoms are the electron density, ρ b, and the Laplacian of electron density, ∇ 2ρ b, at the BCP. In general, a large value of ρ b (>0.2 au) and large and negative value of ∇ 2ρ b indicate a covalent or openshell interaction, whereas a small value of ρ b ( 0 at the BCP generally refers to a closed-shell interaction, i.e., ionic,

van der Waals, or hydrogen bonding. Another useful parameter for characterizing the metal−metal interaction is delocalization index, δ(A,B), defined as the number of electron pairs delocalized between two atoms.22,33 It is generally used as a measure of covalent bond order.33 Due to the reduced covalency, polar bonds have smaller values of δ(A,B) than the predicted bond order.33 Since the metal orbitals involved in bonding are diffuse in nature, many topological properties such as ρ b, ∇ 2ρ b, and even H(r) are subject to change. 33 Nevertheless, the delocalization index, δ(A,B), is almost invariable and produces the bond order expected from simple MO theory.33 Recently, we and others have successfully used the results obtained from QTAIM analysis for an unambiguous description of the nature of intramolecular interactions present in a number of systems.32 For analyzing the nature of metal−metal interactions in group 4−group 9 heterobimetallic complexes, we have performed QTAIM calculations on [Cp2Zr(μ-E)2Rh(H)(CO)L] as a representative case (Table 3). The presence of a BCP Table 3. Electron Density, ρ, Laplacian of Electron Density, ∇ 2ρ, Relative Kinetic Energy Density, G/ρ, and Local Energy Density, H(r), at the Zr···Rh Bond Critical Point of Zr−Rh Heterobimetallicsa structure

ρ

∇ 2ρ

G/ρ

H(r)

4a 4b 4c 4d 5a 5b 5c 5d 6a 6b 6c 6d

0.038 0.027 0.031 0.029 0.036 0.028 0.031 0.032 0.036 0.030 0.033 0.032

0.048 0.038 0.036 0.034 0.047 0.039 0.035 0.035 0.047 0.041 0.038 0.037

0.524 0.470 0.452 0.448 0.503 0.461 0.439 0.428 0.489 0.453 0.436 0.422

−0.008 −0.003 −0.005 −0.004 −0.007 −0.003 −0.005 −0.005 −0.007 −0.007 −0.006 −0.006

a

All values are in au.

between the two metal centers and the formation of two ring critical points (RCP) in the four-membered M−E−M′−E ring in all the representative complexes clearly indicates the presence of a metal−metal interaction in these molecules (Figure 6). Also, the formation of an interatomic surface at the BCP between the two metals further reinforces a bonded interaction.22 We have obtained very small but negative values of H(r) and small and positive values of ρ(r ) and ∇ 2ρ for the above series of complexes at the BCP between the two metal centers (Table 3). This negative value of energy density H(r) indicates a covalent interaction, although the degree may not be very high. It is very clear from Table 3 that there is no significant change in the values for H(r). For a particular bridging group (E), the ρ(r) values increase with variation of L in the order CO < NHC ≈ aNHC < PH3. Although ρ(r) values are not a decisive factor in predicting the nature of a bond, the trend is in agreement with the order of Zr−Rh bond strengths; that is, the shortest Zr−Rh bond has the largest value of ρ(r) at the BCP (Figure 7). This trend is followed for all three bridging groups. When the bridging groups are changed keeping the ancillary ligand (L) fixed, the values of ρ(r ) and ∇ 2ρ do not change appreciably. 5997

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Table 4. Electron Density, ρ, Laplacian of Electron Density, ∇ 2ρ, Relative Kinetic Energy Density, G/ρ, and Local Energy Density, H(r), at the Mo···Rh Bond Critical Point of Mo−Rh Heterobimetallicsa

Figure 6. The trajectory field in the M−E−M′ plane for a group 4− group 9 heterobimetallic complex (4a). Bond paths and basin paths are indicated by black and light blue lines, while the interatomic surfaces are indicated by dark blue lines. Green and red dots indicate bond critical points and ring critical points, respectively.

structure

ρ

∇ 2ρ

G/ρ

H(r)

4a′ 4b′ 4c′ 4d′ 5a′ 5b′ 5c′ 5d′ 6a′ 6b′ 6c′ 6d′

0.040 0.037 0.036 0.035 0.038 0.035 0.033 0.032 0.037 0.035 0.034 0.033

0.043 0.038 0.036 0.035 0.043 0.039 0.035 0.035 0.046 0.040 0.037 0.037

0.500 0.486 0.472 0.457 0.500 0.471 0.460 0.460 0.500 0.474 0.465 0.467

−0.009 −0.008 −0.008 −0.008 −0.007 −0.007 −0.007 −0.007 −0.007 −0.007 −0.006 −0.006

a

All values are in au.

Figure 8. Trajectory field in the M−E−M′ plane for a group 6−group 9 heterobimetallic complex (4a′). Bond paths and basin paths are indicated by black and light blue lines, while the interatomic surfaces are indicated by dark blue lines. Green and red dots indicate bond critical points and ring critical points, respectively.

Figure 7. Variation of M−M′ distances (Å) with electron density (ρ b, au) and its laplacian (∇ 2ρ b, au) for Zr/Rh heterobimetallics (E = SH).

Another parameter that is used to assign the nature of bonding in transition metal complexes is the relative kinetic energy density, Gb/ρ b.33 A value of Gb/ρ b < 1 at the BCP indicates a covalent interaction, whereas Gb/ρ b > 1 indicates the ionic nature of the bond. In our calculations, we have found Gb/ρ b values much smaller than 1.0 for all the complexes (structures 4, 5, 6), which is another indication of the presence of covalency in the metal−metal bonds under consideration. The significant value of the delocalization index δ(Zr, Rh) for structures 4a (0.488) and 5a (0.488) also strengthens this view.33 For the QTAIM analysis of group 6−group 9 heterobimetallic complexes, we have chosen the [(CO)4Mo(μE)2Rh(H)(CO)L] series of complexes as a representative one (Table 4). In these complexes too, we have observed the formation of two ring critical points in the four-membered ring (M−E−M′−E) and a BCP along the bond path between the two metals, which clearly indicates the presence of an M···M′ interaction. Like group 4−group 9 heterobimetallics, here also, the formation of an interatomic surface between two metal atoms further confirms the existence of a metal−metal interaction (Figure 8). At the BCP between the two metals, ρ b(r ) and ∇ 2ρ b(r) values are found to be very small and

positive. The values of ρ b(r) and ∇ 2ρ b(r) nicely correlate with the Mo···Rh bond lengths (Figure 9), with the stronger bonds

Figure 9. Variation of M−M′ distances (Å) with electron density (ρ b, au) and its Laplacian (∇ 2ρ b, au) for Mo/Rh heterobimetallics (E = SH). 5998

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Figure 10. Laplacian distribution along the M···M′ bond path for (a) 4a and (b) 4a′.

having higher values of ρ b(r ) and ∇ 2ρ b(r). The negative value of the total energy density H(r) and small relative kinetic energy density value (Gb/ρ b 0.5 representing higher electron localization. A value of η(r) = 0.5 represents electron gas-like probability. The ELF basins are closely related to Gillespie’s electronic domains and recover the ideas of Lewis. There are two types of basins: (1) core basins denoted as C(A) encompassing the nucleus A and (2) valence basins denoted as V(A, B, ...). The valence basins are characterized by synaptic order,37 which represents the number of core basins sharing a common boundary. Thus, a monosynaptic basin, e.g., V(A), represents lone pairs, while a polysynaptic basin, e.g., V(A, B, ...), represents bi- or polycentric bonds. The presence of valence basin with synaptic order >1 is indicative of shared interaction (covalent, dative, or metallic bonds), while its absence indicates a closed-shell interaction (ionic, van der Waals, or hydrogen bonds).38 The population of electrons in a basin is obtained by integration of the one-electron density of the basin. Figure 11 shows the localization domains of 4a (Figure 11a,b) and 4a′ (Figure 11c,d) at η = 0.33 and 0.30, repectively. The ELF value of 0.33 or 0.30 is rather small, which arises from high contribution of d orbitals in the metal−metal bonding.35b The larger contribution of d orbitals in the metal−metal bond was also evident from the bonding analysis performed with the NBO scheme. Thus, ELF analysis of 4a and 4a′ clearly indicates the presence of direct metal−metal bonds in these heterobimetallic complexes. It is also evident from Figure 10b that a π-like distribution between the metal centers is observed for 4a, indicating a stronger Zr−Rh bond, while a rotationally symmetric localization is found between the metal centers for 4a′. Thus, the ELF analysis also provides evidence for the stronger metal− metal bonds in group 4−group 9 heterobimetallics compared to group 6−group 9 ones. Another topological tool used to describe bonding situations is the source function.39 Bader and Gatti showed that the electron density ρ(r) at any arbitrary point r in a system may be considered to be consisting of contributions from a local source, LS(r, r′), which operates at all other points r′ of space and is given by40

where ∇ 2ρ(r′) is the Laplacian of the electron density at r′, which acts as a source for the electron density at r with an efficiency given by Green’s function |r − r′|−1.39 The density at an arbitrary point r, ρ(r), can be equated as the sum of atomic contributions, S(r, Ω), by integrating LS(r, r′) over the

σ

where D (r) is the excess local kinetic energy due to the Pauli repulsion and Doσ(r) is the kinetic energy of the electron gas having the same density.35b,36 Values of η(r) give a measure of electron localization, i.e., probability of finding electrons alone 5999

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Figure 11. ELF (η) = 0.33 isosurface of (a) 4a and (b) zoomed into the Zr−Mo bonding basin and (c) ELF (η) = 0.30 isosurface of 4a′ and (d) zoomed into the Mo−Rh bonding basin.

Figure 12. Percentage atomic source contributions to the electron density at (a) Zr−Rh BCP (4a) and (b) Mo−Rh BCP (4a′). The positions of the reference points are indicated by black spheres. The volumes of the atomic spheres are proportional to the source contributions from the respective atomic basins.

atom is not much less (12.8%). Figure 12a further reveals that the percentage source contribution of the bridging S atoms is also very significant (10.7%). As a whole, both Zr and Rh as well as the two bridging S atoms contribute to the electron density at the Zr−Rh BCP, indicating a moderately delocalized Zr···Rh interaction. The percentage source function contribution for 4a′ using the Mo−Rh BCP as the reference point (Figure 12b) reveals that the major contribution is coming from the Rh metal (12.0%), while Mo contributes much less (5.7%). It is interesting to note that the bridging sulfur atoms contribute significantly (11.2%), implying that the source is delocalized. The difference in the percentage source contribution of Mo and Rh (6.3%) is higher than that of Zr and Rh (4.0%), suggesting higher polarity in the former. However, contributions of both the metal atoms are significant, and the presence of a bond path and significant values of the delocalization index [δ(Zr,Rh) = 0.488 and δ(Mo,Rh) = 0.442] indicate that in both the heterobimetallics the metal centers are directly bonded.

regions of space bound by the zero flux surfaces of the atomic basins Ω,22

The integrated form of the source function given by S(r, Ω) ≡ ∫ Ω LS(r, r′) dr′, provides a measure of each atom’s contribution to the density at a specific reference point r. This unique topological decomposition provides chemical insights involved in complex bonding situations.41 Figure 12 shows the SF contribution using the BCP of the metal−metal bond in 4a and 4a′ as the reference point. The source function analysis reveals an interesting topological distribution in these supported heterobimetallics. At the Zr−Rh BCP of 4a, it is found that the largest contribution is coming from the Zr atom (16.8%), while the contribution of the Rh 6000

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(8) Rousseau, R.; Stephan, D. W. Organometallics 1991, 10, 3399− 3403. (9) Wark, T. A.; Stephan, D. W. Inorg. Chem. 1990, 29, 1731−1736. (10) Jansen, G.; Schubart, M.; Findeis, B.; Gade, L. H.; Scowen, I. J.; McPartlin, M. J. Am. Chem. Soc. 1998, 120, 7239−7251. (11) Arashiba, K.; Iizuka, H.; Matsukawa, S.; Kuwata, S.; Tanabe, Y.; Iwasaki, M.; Ishii, Y. Inorg. Chem. 2008, 47, 4264−4274. (12) (a) Greenwood, B. P.; Rowe, G. T.; Chen, C. H.; Foxman, B. M.; Thomas, C. M. J. Am. Chem. Soc. 2010, 132, 44−45. (b) Thomas, C. M.; Napoline, J. W.; Rowe, G. T.; Foxman, B. M. Chem. Commun. 2010, 46, 5790−5792. (c) Greenwood, B. P.; Forman, S. I.; Rowe, G. T.; Chen, C. H.; Foxman, B. M.; Thomas, C. M. Inorg. Chem. 2009, 48, 6251−6260. (13) Cooper, B. G.; Fafard, C. M.; Foxman, B. M.; Thomas, C. M. Organometallics 2010, 29, 5179−5186. (14) Krapp, A.; Frenking, G. Theor. Chem. Acc. 2010, 127, 141−148. (15) Takagi, N.; Krapp, A.; Frenking, G. Inorg. Chem. 2011, 50, 819− 826. (16) (a) Finke, R. G.; Gaughan, G.; Pierpont, C.; Cass, M. E. J. Am. Chem. Soc. 1981, 103, 1394−1399. (b) Finke, R. G.; Gaughan, G. Organometallics 1983, 2, 1481−1483. (c) Breen, M. J.; Shulman, P. M.; Geoffroy, G. L.; Rheingold, A. L.; Fultz, W. C. Organometallics 1984, 3, 782−793. (17) (a) Bourissou, D.; Guerret, O.; Gabbai, F.; Bertrand, G. Chem. Rev. 2000, 100, 39−91. (b) Herrmann, W. A. Angew. Chem., Int. Ed. 2002, 41, 1290−1309. (c) Arnold, P. L.; Pearson, S. Coord. Chem. Rev. 2007, 251, 596−609. (d) Albrecht, M. Chem. Commun. 2008, 3601− 3610. (18) B3LYP is Becke’s three-parameter hybrid method using the LYP correlation functional. (a) Becke, A. D. J. Chem. Phys. 1993, 98, 5648− 5652. (b) Lee, C.; Yang, W.; Parr, R. G. Phys. Rev. B 1988, 37, 785− 789. (c) Vosko, S. H.; Wilk, L.; Nusair, M. Can. J. Phys. 1980, 58, 1200−1211. (19) (a) Hay, P. J.; Wadt, W. R. J. Chem. Phys. 1985, 82, 270−283. (b) Wadt, W. R.; Hay, P. J. J. Chem. Phys. 1985, 82, 284−298. (c) Hay, P. J.; Wadt, W. R. J. Chem. Phys. 1985, 82, 299−310. (20) (a) Reed, A. E.; Weinhold, F. J. Chem. Phys. 1983, 78, 4066− 4073. (b) Reed, A. E.; Weinstock, R. B.; Weinhold, F. J. Chem. Phys. 1985, 83, 735−647. (c) Reed, A. E.; Curtiss, L. A.; Weinhold, F. Chem. Rev. 1988, 88, 899−926. (21) Frisch, M. C.; et al. Gaussian 03, revision D.01; Gaussian, Inc.: Wallingford, CT, 2004. For the complete reference, see the Supporting Information. (22) (a) Bader, R. F. W. Atoms in Molecules: A Quantum Theory; Oxford University Press: Oxford, U.K., 1990. (b) Bader, R. F. W. J. Phys. Chem. A 1998, 102, 7314−7323. (c) Bader, R. F. W. Chem. Rev. 1991, 91, 893−928. (23) Keith, T. A. AIMALL (Version 10.02.09); 2010 (http://aim. tkgristmill.com). (24) Godbout, N.; Salahub, D. R.; Andzelm, J.; Wimmer, E. Can. J. Chem. 1992, 70, 560−571. (25) Noury, S.; Krokidis, X.; Fuster, F.; Silvi, B. Comput. Chem. 1999, 23, 597−604. (26) Gatti, C. SF_ESI codes, private communication. (27) Stephan, D. W. Coord. Chem. Rev. 1989, 95, 41−107. (28) (a) Cotton, F. A.; Murillo, C. A.; Walton, R. A. Multiple Bonds between Metal Atoms, 3rd. ed.; Springer: New York, 2005. (b) Pauling, L. The Nature of the Chemical Bond, 3rd. ed.; Cornell University Press: Ithaca, NY, 1960. (29) (a) Guha, A. K.; Sarmah, S.; Phukan, A. K. Dalton Trans. 2010, 39, 7374−7383. (b) Jacobson, H.; Correa, A.; Costabile, C.; Cavallo, L. J. Organomet. Chem. 2006, 691, 4350−4358. (30) Farrugia, L. J.; Evans, C.; Tegel, M. J. Phys. Chem. A 2006, 110, 7952−7961. (31) Cremer, D.; Kraka, E. Angew. Chem., Int. Ed. 1984, 23, 627−628. (32) (a) Phukan, A. K.; Guha, A. K. Inorg. Chem. 2011, 50, 1361− 1367. (b) Phukan, A. K.; Guha, A. K. Inorg. Chem. 2010, 49, 9884− 9890. (c) Flierler, U.; Burzler, M.; Leuser, D.; Henn, J.; Ott, H.; Braunschweig, H.; Stalke, D. Angew. Chem., Int. Ed. 2008, 47, 4321−

4. CONCLUSIONS Density functional calculations at the B3LYP/LANL2DZ level of theory are performed on supported group 4−group 9 and group 6−group 9 heterobimetallic complexes. All these complexes are found to have a significant amount of metal− metal interactions, the strength of which is found to be a function of both the bridging group (E) and the ancillary ligands (L) attached to the late transition metal. Pure σ-donating ancillary ligands such as PH3 lead to a stronger M···M′ interaction, while good and moderate π-acceptor ligands (CO, NHC, aNHC) weaken the same. The strength of the M···M′ interaction is found to be higher in group 4−group 9 than group 6−group 9 heterobimetallic complexes. The chalcogenide-bridged heterobimetallics show a dynamic equilibrium between the cisoid and transoid forms at 298 K and 1 atm pressure. QTAIM analysis reveals that the values of electron density, ρ b(r), and its Laplacian, ∇ 2ρ b(r), at the M···M′ BCPs are very small and positive. However, the values of local electronic densities, H(r), are all negative (albeit smaller) at the M···M′ BCPs. The values of relative kinetic energy density, Gb/ ρ b, and significant values of the delocalization index suggest that the M···M′ interactions present in these molecules are covalent (albeit polar) rather than ionic. ELF analysis further provides evidence for the direct metal−metal bonding by identifying a bonding basin between the two metal centers, while the SF analysis reveals that these metal−metal bonds are moderately delocalized in nature.

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ASSOCIATED CONTENT * Supporting Information Complete reference for Gaussian 03 (ref 21), Tables S1−S5, and Cartesian coordinates of the optimized geometries of all the group 4−group 9 and group 6−group 9 heterobimetallic complexes considered in this study. This material is available free of charge via the Internet at http://pubs.acs.org. S

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AUTHOR INFORMATION

Corresponding Author *Phone: +91 (3712) 267173 (O). Fax: +91 (3712) 267005. E-mail: [email protected].

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ACKNOWLEDGMENTS A.K.P. thanks the Department of Science and Technology (DST), New Delhi, for providing financial assistance in the form of a project (project no. DST/FTP/CS-85/2005). U.G. and A.K.G. thank Tezpur University for an institutional fellowship. We thank all the reviewers for their helpful comments.

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REFERENCES

(1) Wheatley, N.; Kalck, P. Chem. Rev. 1999, 99, 3379−3420. (2) Ritleng, V.; Chetcuti, M. J. Chem. Rev. 2007, 107, 797−858. (3) (a) Gelmini, L.; Stephan, D. W. Organometallics 1988, 7, 849− 855. (b) Cornelissen, C.; Erker, G.; Kehr, G.; Frö hlich, R. Organometallics 2005, 24, 214−225. (4) Hostetler, M. J.; Bergman, R. G. J. Am. Chem. Soc. 1990, 112, 8621−8623. (5) He, Z.; Lugan, N.; Neibecker, D.; Mathieu, R.; Bonnet, J. J. J. Organomet. Chem. 1992, 426, 247−259. (6) Lindenberg, F.; Shribman, T.; Sieler, J.; Hay-Hawkins, E.; Eisen, M. J. Organomet. Chem. 1996, 515, 19−25. (7) Hostetler, M. J.; Butts, M. D.; Bergman, R. G. Organometallics 1993, 12, 65−75. 6001

dx.doi.org/10.1021/om200833a | Organometallics 2011, 30, 5991−6002

Organometallics

Article

4325. (d) Soran, A. P.; Silvestru, C.; Breunig, H. J.; Balázs, G.; Green, J. C. Organometallics 2007, 26, 1196−1203. (e) Cukrowski, I.; Govender, K. K. Inorg. Chem. 2010, 49, 6931−6941. (33) Macchi, P.; Sironi, A. Coord. Chem. Rev. 2003, 238−239, 383− 412. (34) (a) Matito, E.; Solá, M. Coord. Chem. Rev. 2009, 253, 647−665. (b) Gatti, C.; Lasi, D. Faraday Discuss. 2007, 135, 55−78. (35) (a) Becke, A. D.; Edgecombe, K. E. J. Chem. Phys. 1990, 92, 5397−5403. (b) Savin, A.; Nesper, R.; Wengert, S.; Fässler, T. F. Angew. Chem., Int. Ed. Engl. 1997, 36, 1809−1832. (36) Kohout, M.; Savin, A. J. Comput. Chem. 1997, 18, 1431−1439. (37) Silvi, B. J. Mol. Struct. 2002, 614, 3−10. (38) (a) Boily, J. F. J. Phys. Chem. A 2003, 107, 4276−4285. (b) Fowe, E. P.; Therrien, B.; Fink, G. S.; Dual, C. Inorg. Chem. 2008, 47, 42−48. (39) Bader, R. F. W.; Gatti, C. Chem. Phys. Lett. 1998, 287, 233−238. (40) Gatti, C. Struct. Bonding (Berlin) 2011, 1−93, DOI: 10.1007/ 430_2010_31. (41) (a) Gatti, C.; Cargnoni, F.; Bertini, L. J. Comput. Chem. 2003, 24, 422−436. (b) Bertini, L.; Cargnoni, F.; Gatti, C. Theor. Chem. Acc. 2007, 117, 847−884.

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