Density Functional Study of Complexes between Lewis Acids and

Comparative G2(MP2) Study of H3NBX3 and H3PBX3 (X = H, F, and Cl) Donor−Acceptor Complexes. Hafid Anane and Abderrahim Boutalib , Ignacio Nebot-Gil ...
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J. Phys. Chem. 1995,99, 6472-6476

6472

Density Functional Study of Complexes between Lewis Acids and Bases Vicenq Branchadell," Abdelouahid Sbai, and Antonio Oliva Departament de Quimica, Universitat Autbnoma de Barcelona, Edifici Cn, 08193 Bellaterra, Spain Received: November 16, 1994; In Final Form: February 9, 1 9 9 9

Donor-acceptor complexes of ammonia and formaldehyde with M X 3 (M = B, A1 and X = H, F, C1) have been studied using density functional methods. The optimized geometries and the formation energies of the complexes have been compared with ab initio results. The interaction between Lewis acids and bases has been analyzed in terms of steric and electronic effects. A different Lewis acidity scale is obtained depending on the Lewis base. This result has been rationalized in terms of the energy necessary to distort the Lewis acid molecules.

Introduction Donor-acceptor complexes between Lewis acids and bases play a very important role in many catalytic reactions. The knowledge of the structure and properties of these complexes is a necessary goal to understand the mechanism of such processes. The carbonyl group is among the most important functional groups in organic chemistry, so that complexation of carbonyl compounds by Lewis acids is a common step in many organic reactions.' The only experimental data on the geometries of these complexes correspond to solid state structures.2 Several theoretical studies have been devoted to complexes between formaldehyde and boron and aluminum trihalide~.~-' These studies show that reliable results on formation energies can only be obtained at correlated levels and using basis sets with polarization functions. Many theoretical studies have been devoted to complexes involving ammonia as a Lewis base.8-'8 Gas phase structures are known for ammonia-boraneI9 and for trimethylaminealaneZ0 complexes. High level ab initio calculations yield geometries in very good agreement with the experimental data. Brinck et aL2I have recently published an ab initio study of complexes between ammonia and boron trihalides. Their results show that boron trichloride is a stronger Lewis acid than boron trifluoride, in good agreement with the usual Lewis acidity scale. This result has been confirmed by a niore recent study by Jonas et a1.22 This ordering differs from that obtained for aluminum trihalides in =-AI& (X = F, C1) complexes.23 In the last few years, density functional theory (DFT)24-26 has emerged as an altemative to traditional ab initio methods. In this paper DFT will be applied to the study of complexes between Lewis acids and bases. We will focus our attention on complexes of ammonia and formaldehyde with boron and aluminum trihalides, as well as borane and alane, for which ab initio results have been reported. This will allow us to compare the performance of DFT in comparison to ab initio methods. We will also analyze the interaction between both fragments and discuss the relative Lewis acidity scale of the group IIIa trihalides and trihydrides.

Computational Methods All the calculations reported here have been done with the ADF p r ~ g r a m . ~The ~ - geometries ~~ and corresponding energies have been calculated at two levels of theory. At the lowest

* e-mail: [email protected]. @Abstractpublished in Advance ACS Abstracts, April 1, 1995.

level, the local density approximation (LDAh30 with the parameterization of Vosko, Wilk, and N u ~ a i r , has ~ ' been used. At the highest level, LDA has been augmented with the inclusion of the gradient based nonlocal corrections to the correlation and exchange potentials due to P e r d e and ~ ~ B~ e ~ k erespectively. ,~~ The corrections have been included in the SCF procedure34and we will denote these calculations as NL-SCF. For comparison, single point energies based on the LDA geometries have also been estimated by introducing the nonlocal corrections into the LDA energies as a perturbation. We will denote these calculations as NL-P. An uncontracted triple-f STO basis set35has been used for all the atoms. This basis set has been augmented by a set of 3d polarization functions for all nonhydrogen atoms and by a set of 2p polarization functions for H. The frozen-core approximationz8 has been used for all nonhydrogen atoms. A set of auxiliary s, p, d, f, and g STO functions,36centered on all nuclei, has been used to fit the molecular density and represent the Coulomb and exchange potential in each SCF cycle. The numerical integration procedure has been developed by te Velde et The geometry has been fully optimized based on the analytical gradient method developed by Versluis and Ziegler.38 Vibrational frequencies have been calculated at the LDA level in order to compute zero-point vibrational energies through numerical differentiation of energy gradients.39 The formation energy of the complexes has been broken down using the extended transition state (ETS) scheme.40 According to this scheme, the bonding energy between two molecular fragments can be divided into three terms: ~

1

.

~

~

9

~

'

AEsteric is the steric interaction energy and includes the electrostatic interaction energy between the two fragments and the, Pauli exchange repulsion. AEort,ital is due to the interaction between occupied and empty orbitals of the two fragments. Finally, AEpreparationcorresponds to the energy necessary to distort the fragments from their equilibrium geometry to the geometry they have in the complex. These calculations have been done at the NL-P level.

Results and Discussion Figure 1 schematically presents the structures of the complexes NH3-MX3 and HzCO-MX~. In the f i s t case we have only considered the staggered conformation, since it has been shown to be the most table.^^^^^ For the formaldehyde

0022-3654/95/2099-6472$09.00/0 0 1995 American Chemical Society

J. Phys. Chem., Vol. 99, No. 17, 1995 6473

Lewis AcidBase Complexes: Density Functional Study

TABLE 2: Geometw of NH3-BX3 Complexes MX3 levelofcalcn M-N M-X BH3 LDA 1.606 1.219 NL-SCF 1.656 1.221 MP3b 1.664 1.211

Y

exp' LDA

BF3

NL-SCF MP2d expe

a

BC13 A1H3

LDA

NL-SCF MP2d LDA NL-SCF CCSDf expg

AlF3

LDA

AlC13

LDA

NL-SCF NL-SCF

b

1.658 1.642 1.703 1.667 1.60 1.582 1.620 1.618 2.0 15 2.066 2.077 2.063 1.936 1.997 1.952 1.996

1.216 1.368 1.382 1.377 1.36 1.824 1.842 1.839 1.599 1.601 1.585 1.56 1.622 1.646 2.084 2.097

LXMX 113.0 113.5 113.9 113.8 114.1 114.3 114.2 111 113.5 113.5 113.5 117.7 117.5 117.5 114.3 116.9 116.7 116.4 116.1

Figure 1. Schematic representation of the structure of the complexes ammonia-MX3 (a) and formaldehyde-MX3 (b).

Bond lengths in A, bond angles in degrees. Reference 8. Reference 19. Reference 21. e Crystal structure, ref 48. /Crystal structure, ref 15. Trimethylamine-alane complex, ref 20.

TABLE 1: Geometw of MX3 Molecules MX3 level of calcn

TABLE 3: Geometry" of HZCO-MX~ Complexes

BH3

LDA

BF3

NL-SCF LDA NL-SCF expb

BC13

LDA NL-SCF exp'

AlH3

LDA

AIF3 AlC13

NL-SCF LDA NL-SCF expd LDA

NL-SCF expe e

M-X 1.202 1.202 1.308 1.324 1.310 1.736 1.749 1.72 1.583 1.583 1.594 1.616 1.631 2.039 2.049 2.06

Bond lengths in A. Reference 41. Reference 42. Reference 43. Reference 44.

complexes we have also considered the most stable structure according to previous theoretical s t u d i e ~ . ~ - ~ Table 1 presents the values of the optimized geometry parameters corresponding to the isolated Lewis acids. Comparison with the experimental values shows a very good agreement. In general, we can observe that LDA tends to overestimate the strength of the M-X bonds, leading to short bonds. The inclusion of nonlocal corrections increases the bond lengths except for borane and alane, where the LDA and the NL-SCF levels lead to the same results. For these two molecules there are no experimental values available. The results presented in Table 1 can be compared with those obtained for BH3 at the MP4 level, 1.118 A, and by by Pople et Marsh et ~ 1 . for ' ~ at the CCSD level, 1.568 A. Very have reported a DFT study of boron recently, Barone et and aluminum hydrides. For borane and alane they obtain bond lengths of 1.203 A and 1.590 A, respectively. Another DFT study of alane has been recently reported by Pullumbi et aZ.,'" the value obtained for the AI-H bond length being 1.589 A. All these results show a quite good agreement between DFT and high level ab initio results. Tables 2 and 3 present the values of the most important geometry parameters obtained in the geometry optimization of the complexes with ammonia and formaldehyde, respectively.

a

MX3

level of calcn M - 0 C-0

BH3

LDA

NL-SCF MP2b

BF3 LDA NL-SCF HFb BC13 LDA NL-SCF HFC AIH3 LDA NL-SCF HFb A I S LDA NL-SCF AICI3 LDA NL-SCF

HFd

1.511 1.601 1.686 1.702 1.871 2.210 1.589 1.671 1.651 1.907 1.979 2.046 1.878 1.932 1.881 1.951 1.952

1.232 1.237 1.233 1.220 1.226 1.191 1.230 1.237 1.205 1.224 1.232 1.199 1.223 1.232 1.225 1.233 1.215

M-XI M-X2 LMOC LOMXl LOMX2 1.220 1.220 1.209 1.361 1.363 1.318 1.819 1.832 1.832 1.607 1.608 1.609 1.631 1.651 2.094 2.103 2.138

1.220 1.217 1.201 1.342 1.349 1.308 1.802 1.813 1.810 1.593 1.595 1.594 1.613 1.636 2.069 2.081 2.115

123.7 123.1 120.6 118.4 119.3 122.1 123.1 123.1 129.0 122.9 123.4 125.6 120.4 122.6 125.3 125.5 132.6

108.0 106.0 102.7 103.4 101.5 93.4 108.4 108.1 106.5 96.4 96.9 94.6 97.5 98.5 100.7 101.7

104.1 102.3 101.2 100.9 98.8 95.4 102.0 101.5 102.7 99.0 98.8 100.4 100.8 100.3 100.4 100.5

a Bond lengths in A, bond angles in degrees. Reference 3. Reference 4. Reference 5.

In both cases we have included other previous theoretical results from the literature for comparison. Let us first analyze the results corresponding to the ammonia complexes. All complexes present a C3" symmetry and correspond to energy minima, since their computed force constant matrix does not have any negative eigenvalue. The formation of the complexes involves an important pyramidalization of the MX3 molecule. As a general trend one can observe that LDA leads to smaller bond lengths than NL-SCF. For borane and alane complexes there are in the literature experimental gas phase structural data and high level ab initio calculations. The results obtained at the NL-SCF level for the first complex are in better agreement with the experimental data than those obtained at the MP3 level.* For NH3-AlH3 the NL-SCF AI-N and AI-H bond lengths are also very close to the experimental values corresponding to the trimethylamine-alane complex.20 For the LHAlH bond angle the NL-SCF optimization leads to the same value as the CCSD cal~ulation,'~ which is about '3 larger than the experimental value. For the NH3-BF3 complex, the experimental data included in Table 2 correspond to the crystal structure, which is not necessarily a good reference for the theoretical results. Legon and Warnefi9 have reported a N-B bond length of 1.59 8, for

6474 J. Phys. Chem., Vol. 99, No. 17, 1995 TABLE 4: Formation Energie of Complexes Computed at Several Levels of Calculation Lewis base Lewis acid LDA NL-P NL-SCFb ab initio NH3 BH3 -46.5 -31.4 -32.1 (-26.8) -34.7' BF3 -31.8 -18.9 -19.5 (-16.0) -23.2d BC13 -35.6 -21.2 -22.1 (-18.2) -28.2d AIH3 -33.7 -23.7 -25.7 (-22.3) -30.2' AIF3 -46.8 -35.4 -37.5 (-34.3) AlC13 -42.3 -31.0 -32.5 (-29.5) HzCO BH3 -34.0 -17.2 -18.2 (-13.8) -14.4' BF3 -16.8 -3.8 -5.2(-3.5) -9.w -2.0 -3.1 (-0.6) -7.1g BC13 -17.1 AIH3 -24.4 -13.5 -15.7 (-13.0) -20.9 AlF3 -34.5 -22.6 -25.4 (-22.9) AIC13 -28.4 -17.7 -19.8 (-17.4) -26.1h In kcavmol, relative to isolated fragments. Values in parentheses include the zero-point correction computed at the LDA level. MP4/ 6-311G*//MP3/6-31G*,ref 8. dMP2/6-31+G(2d,p),ref 21. CCSD/ DZP, ref 15. fMP3/6-31G*//HF/6-31G*,ref 3. g MP2/6-31G*//HF/6ref 5. 31G*. ref 4. MP4/CEP-31G*//HF/CEP-3lG*, (1

this complex from the gas phase microwave spectra. However, this result has been recently questioned by Jonas and Frenking.I8 The comparison between the NL-SCF and the MP2 results shows a good agreement. The NL-SCF calculation predicts a slightly weaker N-B bond than does MP2. The differences between NL-SCF and MP2 become less important for the NH3BCl3 complex. The comparison between the computed values of the N-A1 bond lengths in the AlF3 and AlC13 complexes shows that they follow the same trend observed in the boron complexes. On the other hand, the weakest N-A1 bond corresponds now to the alane complex. We shall now discuss the geometries of the formaldehyde complexes presented in Table 3. All complexes have C, symmetry and correspond to energy minima, as shown by the analysis of the force constant matrix eigenvalues. The formation of these complexes involves a pyramidalization of the Lewis acid molecules. However, the geometry distortion is less important than that observed in the ammonia complexes, as it can be seen from the values of the M-X bond lengths (see Tables 1, 2, and 3). In all cases one can observe that LDA leads to shorter B-0 and A1-0 bonds than NL-SCF. In the BH3 complex, comparison with the MP2 results shows that this level of calculation the interaction between formaldehyde and the Lewis acid is weaker than at the NL-SCF level. For BF3, the HF calculation clearly underestimates the interaction, leading to a too large value for the B - 0 bond length. On the contrary, for BC13 the HF B-0 bond length is smaller than the NL-SCF value. For the aluminum complexes, the HF calculations lead to values of the A1-0 bond length slightly larger than the NL-SCF results. In all cases HF leads to larger values for the LBOC and LAlOC bond angles than both the LDA and NL-SCF calculations. Table 4 presents the computed values for the formation energy of the complexes relative to the isolated fragments. Comparison between the results obtained at different levels of calculation shows that LDA tends to overestimate the interaction energy between Lewis bases and acids. The introduction of gradient corrections at the NL-P and NL-SCF levels leads to rather similar results. These values are in general below those obtained with ab initio methods, the only exception, being BH3. The inclusion of zero-point vibrational corrections does not change the ordering of the computed values. If one compares the formation energies of the complexes of ammonia with the ones of formaldehyde, one can observe that, for each Lewis acid, ammonia complexes are always stronger

Branchadell et al.

TABLE 5: Decomposition of the Formation Energy of the N H 3 - M X 3 and H2CO-MX3 Complex@ Lewis base MX3 A E s t e n c A E ~ L h~E sIt e n c + A E o r b , t a ~ AEpreparation NH3 BH3 30.5 -75.2 -44.7 12.6 BF, 32.7 -74.5 -41.8 22.3 -42.3 20.2 BC13 74.7 -117.0 -28.6 -29.0 3.3 AIH3 -0.4 -31.3 -42.0 4.5 -10.7 -38.2 5.7 AIC13 1.4 -39.6 -29.7 11.5 HiCO BH3 36.9 -66.6 -16.2 11.0 BF3 18.0 -34.2 -19.7 16.6 BC13 59.3 -79.0 -16.5 0.8 AlH3 10.1 -26.6 -28.5 3.9 A1F3 0.3 -28.8 -23.8 4.0 AlC13 10.0 -33.8 Energies in kcal/mol. than the corresponding formaldehyde complexes. Let us now consider the relative strength of the complexes of each base with respect to the different Lewis acids. For ammonia-BX3 complexes all levels of calculation predict that the strength of the complex increases in the order BF3 < BCl3 < BH3. The ordering between BF3 and BCl3 changes when one considers the formaldehyde-BX3 complexes. Regarding the aluminum complexes, both ammonia and formaldehyde lead to the same ordering and AlF3 forms the strongest complex in both cases. In all cases aluminum halides form stronger complexes than boron halides. On the other hand, borane seems a stronger Lewis acid than alane. In order to understand the different behavior of boron and aluminum halides in front of ammonia and formaldehyde, we have analyzed the interaction between the Lewis acid and base in each complex using the ETS method. Table 5 summarizes the results of the ETS decomposition of the formation energy of the ammonia and formaldehyde complexes. We shall first consider the results corresponding to the ammonia complexes. The values of the steric term show an important difference between boron an aluminum halides. This is due to the Pauli repulsion term contribution, arising from the interaction between occupied orbitals of the Lewis acid, which are mainly centered in the halogen atoms, and the occupied orbitals of ammonia. The different values of M-0 and M-X bond lengths make this repulsion smaller in the aluminum complexes than in the boron complexes. Comparison between the values of the steric term corresponding to borane and boron trihalides shows that this term increases in the order BH3 < BF3 BCl3. This ordering is dominated by the variation of the Pauli repulsion term. The big difference between BF3 and BC13 is due to the electrostatic term, which is related to the different electronegativities of F and C1. On the other hand, for the aluminum halides, the weight of the electrostatic term becomes more important in such a way that the steric term is clearly stabilizing for AlF3. The orbital interaction term is mainly due to the interaction between the LUMO and the Lewis acid and the occupied orbitals of ammonia. Its variation along the BX3 series shows that BC13 is the strongest electron acceptor, while BF3 is a slightly weaker acceptor than BH3. The values of the orbital term corresponding to alane and aluminum trihalides follow a different pattem, since the electron acceptor ability increases in the order AlH3 < AlF3 < AlC13. The addition of the steric term and the orbital term produces a permutation in the ordering between BH3 and BC13, and between AlF3 and AlC13. Finally, the values of the preparation term show an important difference between boron and aluminum compounds. The geometries of the BX3 fragments in the complexes involve X-X distances that are smaller than twice

Lewis AcidlSase Complexes: Density Functional Study

J. Phys. Chem., Vol. 99, No. 17, 1995 6475

of the BX3 moiety. 'Most of the distortion energy is due to

x/B-x

Figure 2. Schematic representation of the geometry distortion of BX3 in the formation of a complex with ammonia (C3" distortion) and with formaldehyde (C, distortion).

TABLE 6: Distortion of the BX3 Molecules in the Formation of the Complexes with Ammonia and Formaldehyde BX3

distortion"

BF3

c 3

BC13

c 3I

"

cs C,

strb

ZLXBX'

AEd

17.4 8.9 27.9 21.1

342.9 (4.75) 351.7 (2.3) 340.5 (5.4) 343.6 (4.6)

21.9 10.0 19.8 15.2

See Figure 2. % of stretching distortion, computed from the values of the B-X bond lengths. In degrees. In parentheses the 8 of angular distortion. Distortion energy in kcaUmol. a

the van der Waals radii of the X atoms. This is not the case for the aluminum complexes, since the AI-X bond lengths are larger than the B-X bond length^.^^,^' Let us now consider the results corresponding to the formaldehyde complexes presented in Table 5. The steric term presents the same differences between boron and aluminum compounds as those observed for the ammonia complexes. The values of the orbital term follow the same ordering observed for the ammonia complexes, and the addition of the steric and orbital terms leads to the same changes. Finally, the preparation energies are in all cases smaller than the ones corresponding to the ammonia complexes. The comparison between the preparation terms in Table 5 shows an important difference in the BF3/BCl3 couple. In ammonia complexes, the preparation energy of BF3 is greater than that of BC13. On the contrary, the situation is the opposite in formaldehyde complexes. This difference in the relative preparation energy of both boron trihalides is responsible for the different ordering of the formation energies of the complexes of ammonia and formaldehyde with the BF3/BC13 couple. In order to analyze this different behavior of the BF3/BCl3 couple in ammonia and formaldehyde complexes, let us consider the distortion of BX3 involved in the formation of both complexes. Figure 2 schematically represents the C3vdistortion involved in the formation of ammonia complexes and the C, distortion corresponding to the formation of formaldehyde complexes, and Table 6 shows the variation of the geometry parameters and the energy associated with these processes. There are two kinds of geometry distortion: the lengthening of the B-X bonds and the pyramidization of the BX3 moiety. The extent of this pyramidalization can be measured through the sum of the LXBX bond angles, this sum being 360" in the equilibrium geometry of boron trihalides. Table 6 shows that the distortion involved in the formation of ammonia complexes is more important than that involved in the formation of formaldehyde complexes. In both cases the geometry of BF3 is less distorted than that of BC13. For formaldehyde there is a direct relation between the degree of geometry distortion and the value of the distortion energy. On the contrary, for ammonia the distortion energy of BF3 is greater than that of BCl3. The distortion of the BX3 molecules in the formation of the ammonia complexes can be formally separated into two steps: a D3h stretching o f the B-X bonds and a C3,, pyramidalization

pyramidali~ation.~~ This fact can be explained if a certain n character of the B-X bond is assumed.53 However, the large values of the pyramidalization energies could also be due to the X-X repulsion20,since the X-X distances become smaller than twice the van der Waals radii of F and C1 in the BX3 complexes.50 The negative charge on the fluorine atoms of BF3 is greater than that on chloride atoms of BCl3, due to the different electronegativities of F and C1, so that the P-F repulsion is greater than the Cl-Cl repulsion.

Concluding Remarks We have studied the donor-acceptor complexes of ammonia and formaldehyde with boron and aluminum trihalides using density functional methods. Comparison with previous ab initio calculations shows that density functional methods are useful in the study of such systems. The results obtained show that ammonia is a stronger Lewis base than formaldehyde and that aluminum halides are stronger Lewis acids than boron halides. We have shown that the Lewis acid acidity scale depends on the based used as a counterpart. While the acidity scale of aluminum halides is the same when compared from ammonia or from formaldehyde complexes, for boron halides the result is different. The interaction between the Lewis base and acid moieties has been analyzed in terms of steric, donor-acceptor, and preparation terms. Through this analysis the role of angular distortion in boron trihalides has been shown to be relevant in the relative acidities of these compounds. Acknowledgment. This work has been supported by DGICYT (Project PB92-0621). The Centre de Parallelisme de Barcelona (CEPBA) is acknowledged for providing computer time in a CONVEX C3480 computer. References and Notes (1) (a) Heathcock, C. H. Assymetric Synthesis; Morrison, P. P., Ed.; Academic Press: New York, 1984; Vol. 3, p 111. (b) Oppolzer, W. Angew. Chem., Int. Ed. Engl. 1984, 23, 876. (c) Reetz, M. T. Angew. Chem., Int. Ed. Engl. 1984, 23, 556. (d) Mikami, K.; Shimizu, M. Chem. Rev. 1992, 92, 1021. (2) Shambayati, S.; Crowe, N. E.; Schreiber, S. L. Angew. Chem., Int. Ed. Engl. 1990, 29, 256 and references cited therein. (3) LePage, T. J.; Wiberg, K. B. J . Am. Chem. SOC.1988, 110, 6642. (4) Branchadell, V.; Oliva, A. J . Am. Chem. SOC.1991, 113, 4132. (5) Jasien, P. G. J . Phys. Chem. 1992, 96, 9273. (6) Coxon, J. M.; Luibrand, R. T. Tetrahedron 1993, 34, 7093. (7) Laszlo, P.; Teston, M. J . Am. Chem. SOC. 1991, 113, 8750. (8) Binkley, J. S.; Thome, L. R. J . Chem. Phys. 1983, 79, 2932. (9) Hirota, F.; Miyata, K.; Shibata, S . J . Mol. Struct. (THEOCHEM) 1989, 201, 99. (10) Chey, J.; Choe, H. S.; Cook, Y. M.; Jensen, E.; Seida, P. R.; Franci, M. M. Organometallics 1990, 9, 2309. (11) Edwards, A. H.; Jones, K. A. J . Chem. Phys. 1991, 94, 2894. (12) Buhl, M.; Steinke, T.; Schleyer, P. v. R.; Boese, R. Angew. Chem., Int. Ed. Engl. 1991, 30, 1160. (13) McKee, M. J. Phys. Chem. 1992, 96, 5380. (14) R ~ g g e n I. , Chem. Phys. 1992, 162, 271. (15) Marsh, C. M. B.; Hamilton, T. P.; Xie, Y.; Schaeffer, H. H. J . Chem. Phys. 1992, 96, 5310. (16) Jungwirth, P.; Zahradnik, R. J . Mol. Struct. (THEOCHEM) 1993, 283, 317. (17) Sakai, S. Chem. Phys. Lett. 1994, 217, 288. ( 18) Jonas, V.; Frenking, G. J . Chem. Soc., Chem. Commun. 1994, 1489. (19) Thome, L. R.; Suenram, R. D.; Lovas, F. J. J. Chem. Phys. 1983, 78, 167. (20) Almenningen, A.; Gundersen, G.; Haugen, T.; Haaland, A. Acta Chem. Scand. 1972, 26, 3928. (21) Brinck, T.; Murray, J. S . ; Politzer, P. Inorg. Chem. 1993,32,2622. (22) Jonas, V.; Frenking, G.; Reetz, M. T. J . Am. Chem. SOC.1994, 116, 8741. (23) Wilson, M.; Coolidge, M. B.; Mains, G. J. J . Phys. Chem. 1992, 96, 4851. (24) Parr, R. G.; Yang, W. Density Functional Theory of Atoms and Molecules; Oxford University Press: New York, 1989. (25) Ziegler, T. Chem. Rev. 1991, 91, 651.

6476 J. Phys. Chem., Vol. 99, No. 17, 1995 (26) Salahub, D. R.; Castro, M.; Proynov, E.i.Relativistic and Electron Correlation Effects in Molecules and Solids; Malli, G. L., Ed.; Plenum Press: New York, 1994. (27) ADF 1.0.2,Department of Theoretical Chemistry, Vrije Universiteit, Amsterdam. (28) Baerends, E. J.; Ellis, D. E.; Ros, P. Chem. Phys. 1973, 2, 41. (29) te Velde, G.; Baerends, E. J. J. Comput. Phys. 1992, 99, 84. (30) Gunnarsson, 0.;Lundquist, I. Phys. Rev. 1974, BlO, 1319. (31) Vosko, S. H.; Wilk, L.; Nusair, M. Can. J . Phys. 1980, 58, 1200. (32) Perdew, J. P. Phys. Rev. 1986, B13, 8822. (33) Becke, A. Phys. Rev. 1988, A38, 3098. (34) Fan, L.; Ziegler, T. J . Chem. Phys. 1991, 94, 6057. (35) Vernooijs, P.; Snijders, G. J.; Baerends, E. J. Slater Type Basis Functionsfor the Whole Periodic System; Internal Report: Freie Universiteit Amsterdam, The Netherland, 198 1. (36) a j n , K.; Baerends, E. J. Fit Functions in the HFS Methodr; Internal Report: Freie Universiteit Amsterdam, The Netherlands, 1984. (37) Boenitger, P. M.; te Velde, G.; Baerends, E. J. Int. J. Quantum Chem. 1988, 33, 87. (38) Versluis, L.; Ziegler, T. J . Chem. Phys. 1988, 88, 322. (39) Fan, L.; Versluis, L.; Ziegler, T.; Baerends, E. J.; Ravenek, W. Int. J. Quantum Chem. 1988, S22, 173. (40) Ziegler, T.; Rauk, A. Theor. Chim. Acta 1977, 46, 1. (41) Brown, C. W.; Overend, J. Can. J . Phys. 1968, 46, 9077. (42) CRC Handbook of Physics and Chemistry, 69th ed.; Weast, R. C., Ed.; CRC Press: Boca Raton, FL, 1988.

Branchadell et ai. (43) Girichev, G. B.; Utkin, A. N.; Giricheva, N. I. Izv. Vys. Uch. Zav. Khim. Khim. Teckhnol. 1983, 26, 634. (44)Tomita, T.; Sjogren, C. E.; Klaeboe, P.; Papatheodoru, G. N., Rytter, E. J . Raman Spectrosc. 1983, 14, 415. (45) Pople, J. A.; Luke, B. T.; Frisch, M. J.; Binkley, J. S. J . Phys. Chem. 1985, 89, 2198. (46) Barone, V.; Adamo, C.; Flisz&, S.; Russo, N. Chem. Phys. Lett. .> 1994, 222,597. (47) Pullumbi. P.: Bouteiller. Y.: Manceron. L.: Miioule. C. Chem. Phvs. 1994, 18.5, 25. (48) Hoard, J. L.; Geller, S.; Cashin, W. M. Acta Crystallogr. 1951, 4 , 396 _ _ .. (49) Legon, A. C.; Warner, H. E. J. Chem. Soc., Chem. Commun. 1991, 1397. (50) The values of the F-F distances are 2.32 %, for ammonia-BF3 and 2.80 8, for ammonia-AlF3. The values of C1-C1 distances are 3.08 8, for ammonia-BC13 and 3.56 A for ammonia-AlC13. All these values correspond to the NL-SCF calculations. (51) The van der Waals radii of F and C1 ar 1.35 8, and 1.80 A, re~pectively.~" (52) The energy associated with the D3h stretching distortion is 5.4 kcaU mol for BF3 and 4.8 kcaVmol for BCl3. (53) Branchadell, V.; Oliva, A. J . Mol. Struct. (THEOCHEM) 1991, 236, 75.

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