Theoretical investigation of C HH B dihydrogen bonded complexes of acetylenes with borane-trimethylamine

Transcription

1 Theoretical investigation of C HH B dihydrogen bonded complexes of acetylenes with borane-trimethylamine Prashant Chandra Singh, G. Naresh Patwari * Department of Chemistry, Indian Institute of Technology Bombay, Powai, Mumbai , India Abstract Formation of C HH B dihydrogen bonded complexes of acetylene, fluoroacetylene, chloroacetylene, and cyanoacetylene with borane-trimethylamine were investigated with MP2 and B3LYP methods using G(d,p) and aug-cc-pvdz basis sets. The stabilization energies ranged from 6 20 kj mol 1. NBO analysis predicts transfer of charge from r B H bonding orbital to r* C H anti-bonding orbital. It was also found that the lowering of the C H stretching frequencies of acetylene moiety in the dihydrogen bonded complex does not correlate with other hydrogen bonded complexes, indicating that the supposition of dihydrogen bonding as another type of hydrogen bonding (r and p) may be incorrect. 1. Introduction Over the past few years several cases of hydrogen bonds involving unconventional donors and/or acceptors have been reported [1]. For instance, hydrogen bonds in which p bond electrons act as acceptors [2] and also hydrogen bonds involving C H groups as donors [3] have been well documented. Yet another important class of unconventional hydrogen bonds are the interaction of oppositely charged hydrogen atoms, which can be generically represented as E H(d )(d+)h X, wherein E and X are atoms which are less and more electronegative than hydrogen, respectively. This favorable interaction between oppositely charged hydrogens has been termed as Ôdihydrogen bondingõ [4,5]. Apart from metal hydrides, borane-amines have played crucial role in understanding dihydrogen bonding. Notable are the investigation of dihydrogen bonding in BH 3 NH 3 [6 9] and its complexes with molecules containing acidic hydrogens [10 12]. MikamiÕs group has reported the formation of B HH X (X = O, N) dihydrogen bonded complexes in the gas phase [13]. The ability of C H bonds to form hydrogen bonds has been fairly well recognized, albeit weaker than the conventional hydrogen bonds [3]. This information can be extrapolated to infer the formation of dihydrogen bonds involving CH groups. Recently, Grabowski et al. [14] have reported the formation of dihydrogen bonded complexes between acetylene and several hydrides, such as LiH, BeH 2, and HBeF. Further, to the best of our knowledge, reports on C HH B dihydrogen bonds are extremely sparse [15], and thus provides the impetus for investigating of C HH B type dihydrogen bonded systems. In this Letter, we report the formation of dihydrogen bonded complexes of borane-trimethylamine (BTMA) with acetylene and F, Cl, CN substituted acetylenes. Along with the energetic aspects of the interaction, structural and spectroscopic markers were computed using a high level of theory. Comparison with other hydrogen bonded complexes of acetylene is also presented. 2. Methodology Theoretical methods employed in this report are MP2(FC) and DFT-B3LYP, both of which include electron correlation, using G(d,p) and aug-cc-pvdz

2 6 basis sets. It has been shown in several instances that the inclusion of polarization and diffuse functions on heavy atoms improves the description of hydrogen bonding [16]. Since dihydrogen bonding is essential an interaction, analogous to hydrogen bonding, between a pair of oppositely charged hydrogen atoms, polarization and diffuse functions have been added on hydrogen as well. The equilibrium structures of the monomers and the complexes were calculated and the nature of stationary point obtained was confirmed by calculating the vibrational frequency at the same level of theory. The calculated vibrational frequencies were scaled by 0.958, 0.958, 0.951, and 0.96 for B3LYP/ G(d,p), B3LYP/aug-cc-pVDZ, and MP2/ G(d,p), MP2/aug-cc-pVDZ methods, respectively. The scaling factors were adapted to match the experimental C H stretching frequencies of acetylene [17]. According to Kim et al. [18] 100% of BSSE correction often underestimates the interaction energy and 50% correction is a good empirical approximation. The interaction energies, reported herein, are corrected for the scaled zero point vibrational energy (szpve) and 50% BSSE correction. It is well known that Mulliken charges are not adequate to explain the bonding in various situations, therefore molecular electrostatic potential (MEP) derived charges as well as charges from the natural population analysis (NPA) have also been obtained. The topographical analysis of the electron density distribution has been carried using Atom in Molecules approach [19]. Further, to get more insights about the lowering of the C H stretching frequency of the donors upon complex formation, natural bond analysis (NBO) has also been carried out [20]. All the calculations reported here were carried out using GAUSSIAN 98 [21]. 3. Results and discussion The HH distance of less than 2.4 Å, twice the van der Waals radius of the hydrogen atom (1.2 Å), is the most used geometrical criterion to identify the formation of dihydrogen bonds. However, due to electrostatic nature of the (di)hydrogen bonding, van der Waals cutoff criteria is strongly criticized, as electrostatic interaction acts beyond this distance [22]. Moreover, it has been reported recently that for C H bonds, the van der Waals radius of hydrogen atom is marginally greater than 1.2 Å [23]. However, to be inline with the existing reports in the literature, we have considered 2.4 Å as the cutoff of the formation of dihydrogen bonds. Fig. 1 depicts the calculated structure of acetylene-btma complex at MP2/aug-cc-pVDZ level of theory, which has a pair of symmetrically bifurcated dihydrogen bonds, with HH distance of 2.27 Å. Further, the acetylene-btma complex was also investigated at MP2/aug-cc-pVTZ level. In this case the geometry changes marginally with respect to MP2/aug-cc-pVDZ level and the HH distance is slightly elongated by Å. All the other complexes form similar structures and Table 1 lists Fig. 1. MP2/aug-cc-p VDZ calculated structure of acetylene-btma complex. Hydrogen atoms are depicted in grey, symbols B, C, and N correspond to boron, carbon and nitrogen atoms, respectively, and the distances are given in Å. the relevant intermolecular geometrical parameters. Further, for all the complexes the B HH angles are the range of 92 97, while C HH angles, comparatively more linear, are in the range of The present results are in accord with those reported elsewhere [5 13]. Table 1 also lists the stabilization energies corrected for szpve and 50% BSSE for various dihydrogen bonded complexes. It is apparent from Table 1 that for all the complexes the MP2 stabilization energies are higher than the corresponding DFT-B3LYP value. This can be attributed to the neglect of the dispersion energy contribution in the DFT-B3LYP method, which can be significant for complexes involving C H groups. The highest stabilization energies are obtained at MP2/aug-cc-pVDZ level of theory, are in agreement with the shortest distances obtained at the same level. It can also be seen from Table 1 that the stabilization energy increases with the substitution of fluoro, chloro, and cyano groups, in that order, respectively. For instance at MP2/aug-cc-pVDZ level the stabilization energy of acetylene-btma complexes is 14.2 kj mol 1 increases to 20.6 kj mol 1 in the case of cyanoacetylene complex, while the fluoro- and chloro-acetylenes taking intermediate values of 14.8 and 16.2 kj mol 1, respectively. In order to understand the nature of HH interaction in various complexes, we carried out charge analysis using molecular electrostatic potential (MEP) method as well as the natural population analysis (NPA) scheme at MP2/ aug-cc-pvdz level and the results are presented in Table 2. In the case of acetylene-btma complex, the MEP charge on the acetylenic hydrogen is , while the two interacting hydrogens of BH 3 group show a negative charge AU. The NPA results show that the acetylenic hydrogen has a positive charge of AU, while the interacting hydrogens of borane have a negative charge is AU, thus confirming the formation of dihydrogen bond between two oppositely charged hydrogen atoms. The same trend continues for all the complexes with positive charged acetylenic hydrogen interacting with two negative charged hydrogens of BH 3 group forming a pair of symmetrically bifurcated dihydrogen bonds.

3 7 Table 1 Optimized HH distances (Å), B HH and C HH angles ( ) and stabilization energies, DE, (kj mol 1 ) for various dihydrogen bonded complexes of BTMA B3LYP/ G(d,p) B3LYP/aug-cc-pVDZ MP2/ G(d,p) MP2/aug-cc-pVDZ C 2 H 2 BTMA HH B HH C HH DE FC 2 H BTMA HH B HH C HH DE ClC 2 H BTMA HH B HH C HH DE NCC 2 H BTMA HH B HH C HH DE Table 2 MEP and NPA charges on interacting acetylenic and borane hydrogens in all the four complexes calculated at MP2/aug-cc-pVDZ level of theory MEP NPA CH BH CH BH C 2 H 2 BTMA FC 2 H BTMA ClC 2 H BTMA NCC 2 H BTMA The bonding pattern of dihydrogen bonded complexes can be analyzed by the topographical study of electron density. In the present case, the critical points of electron density distribution were obtained, characterized by the rank and trace of the Hessian matrix. A (3, 1) bond critical point with positive Laplacian for the electron density distribution at the bond critical point indicates the non-covalent interaction. However, in the case of very strong hydrogen and dihydrogen bonds, which are often partly covalent in nature, have negative values of Laplacian of electron density at the BCPs [24]. Table 3 lists the electron density and their Laplacian for the present complexes. The electron density and its Laplacian range from to 0.04 and 0.02 to 0.15 a.u., respectively, and are in the range acceptable for hydrogen [25] and dihydrogen bonds [26]. Further, the BCPs between the two interacting hydrogens in all the complexes are of the (3, 1) type, with positive values for the Laplacian of the electron density. It is clearly evident from the charge and critical point analysis that in all the complexes the interaction is between two oppositely charged hydrogen atoms thus confirming the formation of dihydrogen bonds. Further, natural bond orbital (NBO) analysis was also carried out at B3LYP/aug-cc-pVDZ level using structures optimized at MP2/aug-cc-pVDZ level. This is due to the fact that the natural bond orbital (NBO) analysis yields reliable values only with either HF or DFT methods, since the wavefunction is completely defined [27]. Table 3 also lists the changes in the electron occupancies, relative to the monomer, for both donor r B H bonding orbital and the acceptor r* C H bonds anti-bonding orbital. Clearly, the donor orbital looses the occupancy and the acceptor gains, implying charge transfer from r B H bonding orbital to r* C H bonds anti-bonding orbital. Consequently the C H bond is weakened resulting in lowering its stretching frequency. It is well known that the donor X H bond is elongated due to the formation of hydrogen bond. In the case of Table 3 Electron density at the critical points (CP), Laplacian of electron density between HH contacts and the changes in the electron occupancy in the bonding a and anti-bonding b orbitals of the donor and the acceptors, respectively, in various complexes Nature of CP q(r) D 2 q(r) DDonor occupancy DAcceptor occupancy C 2 H 2 BTMA (3, 1) FC 2 H BTMA (3, 1) ClC 2 H BTMA (3, 1) NCC 2 H BTMA (3, 1) a r BH is the donor NBO. b r* CH is the acceptor NBO.

4 8 Table 4 C H stretching frequencies a and their shifts b in various complexes of acetylenes with BTMA B3LYP/ G(d,p) B3LYP/aug-cc-pVDZ MP2/ G(d,p) MP2/aug-cc-pVDZ C 2 H 2 BTMA 3243 (3275) (3274) (3288) (3292) (3374) (3373) (3373) (3375) 18 FC 2 H BTMA 3284 (3345) (3349) (3357) (3368) 80 ClC 2 H BTMA 3273 (3337) (3343) (3342) (3355) 76 NCC 2 H BTMA 3246 (3322) (3328) (3321) (3333) 95 a The values given in parenthesis are for the monomer. b Shifts are italicized. BTMA complexes with the acetylenes, we have examined the changes in the bond lengths and the stretching frequencies of the C H groups. In the case of acetylene-btma complex the maximum elongation of the interacting C H bond (0.004 Å) is observed at MP2/aug-cc-pVDZ level, while the other C H bond of acetylene showed marginal elongation of Å. On the acceptor side the two interacting B H bonds are elongated by Å, and the non-interacting B H bond shortens by same amount. In the case of fluoro-, chloro-, and cyano-acetylene complexes with BTMA, the elongation of C H bonds are 0.006, 0.006, and Å, respectively, at MP2/aug-cc-pVDZ level of theory. The natural consequence of the elongation of the C H bonds is lowering of its stretching frequency. The vibrational spectroscopy of the donor groups in the hydride stretching region has been one of the major experimental tool to characterize hydrogen bonds [1]. Table 4 lists the C H stretching frequencies of acetylenes and the shifts upon dihydrogen bonded complex formation. Once again, in all the cases the lowering of the C H stretching frequency is higher when calculated using aug-cc-pvdz basis set. Also, as expected the lowest shifts are observed for acetylene-btma and highest for cyanoacetylene- BTMA complex, while the fluoro- and chloro-acetylene complexes take intermediate values. The two most important factors that influence the structure and energetics of hydrogen bonding are: (1) the acidity of the donor; (2) the proton affinity of the acceptor. Therefore, for a given donor the stabilization energy should be proportional to the proton affinity of the acceptor. For example, MikamiÕs group has shown for several hydrogen bonded complexes of phenol that the lowering of the O H stretching vibration of the phenol moiety is linearly correlated to the proton affinity of the acceptor [28,29]. To ascertain whether the linear correlation holds for hydrogen bonded complexes of acetylene, the structures and vibrational frequencies of hydrogen bonded complexes of acetylene with water, methanol, dimethylether, ammonia, trimethylamine, and N,N-dimethyl-ethylamine were calculated at MP2/ G(d,p) level and are listed in Table 5. Fig. 2 shows the plot of total lowering of C H stretching frequencies [R(Dm CH )] of acetylene in hydrogen bonded complexes with water, methanol, ammonia, trimethylamine, and N,N-dimethyl-ethylamine against the gas-phase proton affinities of the acceptors. It is evident from Fig. 2 that the lowering of C H stretching frequencies of the acetylene moiety is linearly correlated with proton affinities of the bases. In order to set the confidence limit, using the calculated R(Dm CH )of98cm 1 for the acetylene-dimethylether complex (Table 5), the proton affinity of dimethylether is estimated to be 780 kj mol 1 (Fig. 2, m), which is in very good agreement with the experimental value of kj mol 1 [30]. Further, from the same correlation, using R(Dm CH )of49cm 1 for the dihydrogen bonded acetylene-btma complex, the proton affinity of BTMA can be estimated as 690 kj mol 1 (Fig. 2, n). Recently, we had used the linear correlation between the lowering of the O H stretching vibration of the phenol moiety in various hydrogen bonded complexes and estimated the proton affinity of BTMA as 710 kj mol 1 [31]. Table 5 C H stretching frequencies (in cm 1 ) of acetylene and its complexes calculated at MP2/ G(d,p) level of theory m 1 m 3 Dm 1 Dm 3 R(Dm) PA/kJ mol 1 C 2 H C 2 H 2 BTMA a, b C 2 H 2 C 2 H c C 2 H 2 H 2 O c C 2 H 2 MeOH c C 2 H 2 MeOMe a, c C 2 H 2 NH c C 2 H 2 NMe c C 2 H 2 NEtMe c a Estimated using linear correlation in Fig. 2. b G2 proton affinity from [31]. c Gas phase proton affinities from [30].

5 NMe 3 NEtMe 2 Acknowledgments Financial support from CSIR (Grant No. 01(1902)/03/ EMR-II) and BRNS (Grant No. 2004/37/5/BRNS/398) are gratefully acknowledged. Σ( ν) CH / cm The estimated values of proton affinity of BTMA, in both the cases (690 and 710 kj mol 1 ) are much lower than the calculated value of kj mol 1 using G2 method [31], which is know to predict the proton affinities with an accuracy of ±10 kj mol 1 [32 34]. From these results, one can straightforwardly infer that the linear correlation between frequency shifts and the proton affinities involving hydrogen bonded species grossly underestimates the proton affinities of the dihydrogen bonded species. This implies that such a correlation does not hold for dihydrogen bonded complexes and re-emphasizes our earlier assertion [31] that the premise of the dihydrogen bond being another type of hydrogen bond (r and p) might be incorrect. 4. Conclusions H 2 O C 2 H 2 BTMA MeOMe MeOH NH Proton Affinity / kjmol -1 Fig. 2. Plot of total lowering of the C H stretching frequencies R(Dm) of acetylene moiety in various hydrogen bonded systems vs proton affinities of the acceptors (h) The straight line is linear least-squares fit to the data points, excluding the MeOMe (m) and BTMA (j). From the fit, the proton affinities of MeOMe and BTMA can be estimated as 780 and 690 kj mol 1, respectively. In summary, acetylene and fluoro, chloro, and cyano substituted acetylenes form dihydrogen bonded complex with borane-trimethylamine, with HH contact distances less than 2.4 Å and stabilization energies in the range of 6 20 kj mol 1. The stabilization energy increases with the substitution of fluoro, chloro, and cyano groups, in that order, respectively. The molecular electrostatic potential derived charge, natural population, natural bond order, atoms in molecules analysis support the formation of C HH B dihydrogen bonds. Using the linear correlation between the lowering of the C H stretching frequencies of acetylene and proton affinities of bases, in various hydrogen and dihydrogen bonded complexes, the proton affinity of borane-trimethylamine was estimated as 690 kj mol 1, a value much lower than the G2 value of kj mol 1. It can be concluded from the results presented here that dihydrogen bonded complexes cannot be correlated along with the hydrogen bonded complexes. References [1] For example, see: A. Fujii, G.N. Patwari, T. Ebata, N. Mikami, Int. J. Mass. Spectrom. 202 (2002) 289. [2] S. Suzuki, P.G. Green, R.E. Baumgarner, S. Dasgupta, W.A. Goddard III, G.A. Blake, Science 257 (1992) 942. [3] G.R. Desiraju, Acc. Chem. Res. 29 (1996) 441. [4] R.C. Stevans, R. Bau, R. Milstein, O. Blum, T.F. Koetzle, J. Chem. Soc., Dalton Trans. 4 (1990) [5] L.S. Vander Sluys, J. Eckert, O. Eisenstein, J.H. Hall, J.C. Huffman, S.A. Jackson, T.F. Koetzle, G.J. Kubas, P.J. Vergamini, K.G. Caulton, J. Am. Chem. Soc. 112 (1990) [6] W.T. Klooster, T.F. Koetzle, P.E.M. Siegbahn, T.B. Richardson, R.H. Crabtree, J. Am. Chem. Soc. 121 (1999) [7] G. Merino, V.L. Bakhmutov, A. Vela, J. Phys. Chem. A 106 (2002) [8] C.A. Morrison, M.M. Siddick, Angew. Chem. Int. Ed. 43 (2004) [9] T.B. Richardson, S.D. Gala, R.H. Crabtree, P.E.M. Siegbahn, J. Am. Chem. Soc. 117 (1995) [10] Y. Meng, Z. Zhouc, C. Dunab, B. Wangd, Q. Zhonga, J. Mol. Struct. (Theochem) 713 (2005) 135. [11] S.A. Kulkarni, J. Phys. Chem. A 102 (1998) [12] T. Kar, S. Scheiner, J. Chem. Phys. 119 (2003) [13] G.N. Patwari, T. Ebata, N. Mikami, Chem. Phys. 283 (2002) 193. [14] S.J. Grabowski, W.A. Sokalski, J. Leszcynski, J. Phys. Chem. A 108 (2004) [15] T. Ramnial, H. Jong, I.D. McKenzie, M. Jennings, J.A.C. Clyburne, Chem. Commun. 14 (2003) [16] T.S. Zwier, J. Phys. Chem. A 105 (2001) [17] The experimental symmetric and asymmetric C H stretching frequencies of acetylene are 3374 and 3289 cm 1, respectively, 79 edn.d.r. Lide (Ed.), CRC Handbook of Chemistry and Physics, CRC Press, Boca Raton, [18] K.S. Kim, P. Tarakeshwar, J.Y. Lee, Chem. Rev. 100 (2000) [19] R.F.W. Bader, Atoms in Molecules. A Quantum Theory, Oxford University Press, New York, [20] A.E. Reed, L.A. Curtiss, F. Weinhold, Chem. Rev. 88 (1988) 899. [21] M.J. Frisch et al., GAUSSIAN 98, Revision A.11.4, Gaussian, Inc., Pittsburgh, PA, [22] G.R. Desiraju, T. Steiner, The Weak Hydrogen Bond in Structural Chemistry and Biology, Oxford University Press Inc., New York, [23] B. Lakshmi, A.G. Samuelson, K.V.J. Jose, S.R. Gadre, E. Arunan, New J. Chem. 29 (2005) 371. [24] S.J. Grabowski, J. Phys. Chem. A 105 (2001) [25] U. Koch, P.L.A. Popelier, J. Phys. Chem. A 99 (1995) [26] P.L.A. Popelier, J. Phys. Chem. A 102 (1998) [27] I. Alabugin, T.A. Zeidan, J. Am. Chem. Soc. 124 (2002) [28] A. Iwasaki, A. Fujii, T. Ebata, N. Mikami, J. Phys. Chem. A 100 (1996) [29] A. Fujii, T. Ebata, N. Mikami, J. Phys. Chem. A 106 (2002) [30] E.P.L. Hunter, S.G. Lias, J. Phys. Chem. Ref. Data 27 (1998) 413. [31] G.N. Patwari, J. Phys. Chem. A 109 (2005) [32] L.A. Curtiss, K. Raghavachari, G.W. Trucks, J.A. Pople, J. Chem. Phys. 94 (1991) [33] B.J. Smith, L. Radom, J. Am. Chem. Soc. 115 (1993) [34] A.L.L. East, B.J. Smith, L. Radom, J. Am. Chem. Soc. 118 (1997) 9014.