J . Phys. Chem. 1989, 93, 3929-3936
3929
A Molecular Orbital Study of Azole-Li+ Complexes M. Alcami, 0.M6, and M . Yhfiez* Departamento de Quimica, C-XIV, Facultad de Ciencias, Universidad AutBnoma de Madrid, Canto Blanco, 28049 Madrid, Spain (Received: August 8, 1988: In Final Form: December 2, 1988)
Hartree-Fock calculations with the 6-3 1G* basis have been performed to investigate the structure and Li+ binding energies of the complexes between Li' and a series of azoles. Structures have been fully optimized at the 3-21G level. A topological analysis of the Laplacian of the electronic charge density reveals that the nature of Li+-azole complexes is markedly different from that of protonated azoles. In the former, stabilization arises mainly from electrostatic interactions while in the latter a new covalent bond between the basic nitrogen and the incoming proton is actually formed. Basis set superposition error (BSSE) is quite significant, especially for Lic-bridging complexes. For those cases where Li+ is single coordinated a good linear correlation between Li+ binding energies and proton affinities is found, the former being 4 times smaller than the latter. A similar correlation is obtained for the corresponding relative values, although in this case the attenuation factor is only 1.8, as it should be expected from simple ion-dipole interactions. When the azole presents two nitrogen atoms having a lone pair of electrons, Li+-bridging structures are not always the most stable forms. A discussion on which factors favor this particular conformation is offered.
Introduction The interactions of acids and bases are of great importance in chemistry. Moreover, the possibility of measuring both H+ and Li+ affinities with high accuracy, by means of different experimental technique^,'-^ has stimulated a growing interest in the study of these interactions. In this respect, comparison of proton and Li+ affinities reveals4 that the relative basicities of a given set of bases usually depend on the reference acid; furthermore, substituent effects on the strength of a particular base usually depend on the reference acid too. In our research on the gas-phase basicity of different organic basess-I6 we have lately focused our attention on the behavior of azoles"-'4 because they are strong bases in the gas phase and because they constitute a homogeneous set of compounds where the number of basic centers increases steadily from pyrazole and imidazole to pentazole. This is a quite important fact since although there is a considerable number of studies on complexes containing an ion and monodentate ligands, the studies on complexes of ions and bases which present more than one basic center are scarcer. As a typical example of this kind of studies, the recent work of Houriet et aI.l7 on the gas-phase basicity of @-amino (1) Bowers, M. T.; Aue, D. H.; Webb, H . M.; McIver Jr., R. T. J. Am. Chem. SOC.1971, 93, 4314. (2) Briggs, J. R.; Yamdagni, R.; Kebarle, P. J. Am. Chem. SOC.1972,94, 5128. (3) Woodin, R. L.; Beauchamp, J . L. J. Am. Chem. SOC.1978,100, 501. (4) Woodin, R. L.; Houle, F. A,; Goddard 111, W. A . Chem. Phys. 1976, 14, 461. ( 5 ) Catalln, J.; M6, 0.;Perez, P.; Yliiez, M. J. Am. Chem. SOC.1979, 101, 6520. (6) Catalln, J.; Perez, P.; YlRez, M. Tetrahedron 1982, 38, 6393. ( 7 ) Catalln, J.; Mb, 0.;Perez, P.; YiRez, M . Tetrahedron 1983, 39, 2851. (8) Catalln, J . ; M6, 0.; Perez, P.; YlRez, M . J. Mol. Struct.: THEOCHEM 1983, 94, 143. (9) Catalln, J.; Mb, 0.;Perez, P.; Yliiez, M.; Amat-gueri, F. N o w . J. Chim. 1984, 8. 87. CIO) Catalln, J.; de Paz, J . L. G.; Yliiez, M.; Elguero, J. J. Mol. Struct.: THEOCHEM 1984. 108. 161: J . Am. Chem. SOC.1984. 106, 6552. ( I I ) Catalln, J.; de Paz, J. L. G.;YlRez, M.; Elguero, J. Chem. Scr. 1984, 24, 84. (12) Catalln, J.; Mb, 0.;de Paz, J. L. G.; Perez, P.; Yiiiez, M.; Elguero, J . J . Org. Chem. 1984, 49, 4319. (13) M6,O.; de Paz, J. L. G.; Yliiez, M. J. Phys. Chem. 1986, 90, 5597. (14) Mb, 0.;YlRez, M.; Elguero, J . J. Org. Chem. 1987, 52, 1713. (15) Mb, 0.;de Paz, J. L. G.; YlAez, M. J. Phys. Chem. 1987, 91, 6484. (16) Catalln, J.; de Paz, J . L. G.; Yliiez, M.; Amat-Guerri, F.; Houriet, R.; Rolli, E.; Zehringer, R.; Oelhafen, P.; Taft, R. W.; Anvia, F.;Qian, J . H . J. Am. Chem. SOC.1988, 110, 2699. ( I 7) (a) Houriet, R.; Rufennacht, H.; Carrupt, P. A,; Vogel, P.; Tichy, M. J. Am. Chem. SOC.1983, 105, 3417. (b) Guenat, C.; Houriet, R.; Stahl, D.; Winkler, J. Helv. Chim. Acta 1985, 68, 1647. (c) Bollinger, J . C.; Houriet, R.; Kern, C. W.; Perret, D.; Weber, J.: Yvernault, T. J. Am. Chem. Sot. 1985, 107. 5352.
open-chain and cyclic diols,17band aliphatic phosphine oxides and p h o s p h ~ r a m i d e s should l ~ ~ be mentioned. However, most of the information available for bidentate bases is of theoretical Actually, a b initio studies are a powerful tool in evaluating not only binding energies for complexes involving different ions but also their most stable structures, information which is seldom amenable from experiments. Azoles constitute a particular subset of polydentate bases, since they are heteroaromatic compounds containing only nitrogen as heteroatoms, and therefore they provide a suitable set of systems to study these multicenter interactions without changing the nature of the basic center. In previous work we have already considered the structures and energies of all protonated azolesI3 as well as those of the complexes between methylpyrazoles and methylimidazoles with H+ l 2 and NH4+.I4In this paper we extend this study of the interactions of the same set of bases with the Lif acid. Besides, Li+-azole complexes are also of an intrinsic interest since, although the gas-phase proton affinity of several members of the family has been already measured,25 as far as we know no experimental information is available on their Li' binding energies. On the other hand, they may be used as model systems to dig deeply into the understanding of the coordination chemistry of alkali-metal ions with nucleobases,26and they may be also important in the field of cation-selective transport through biological membra ne^.^^,^^
In the present work Li+-azole binding energies are examined as well as the relative stabilities of the different conformers we can define for each particular compound. In particular, we shall discuss why the bridged structure is not always the most stable one as it has been found for smaller bases. Consequently, a detailed analysis of the nature of the Li'-azole interaction will be carried out in these cases. Of particular interest will be the comparison of the present data for Li' binding energies with the proton affinities obtained previ~usly'~ at the same level of accuracy. (18) Miertus, S.; Kysel, 0. Chem. Phys. Lett. 1975, 35, 531. (19) Del Bene, J. E.; Frisch, M. J.; Raghavachari, K.; Pople, J. A. J . Phys. Chem. 1982,815, 1529. (20) Del Bene, J . E.; Frisch, M. J.; Raghavachari, K.: Pople, J . A,; Schleyer, P. v. R. J. Phys. Chem. 1983, 87, 73. (21) Del Bene, J. E. J . Phys. Chem. 1983, 87, 367. (22) Del Bene, J . E. J. Phys. Chem. 1984, 88, 5927. (23) Ikuta, S. Chem. Phys. Lett. 1985, 116, 482. (24) Ikuta, S. Chem. Phys. 1984, 108, 441. (25) Meot-Ner (Mautner), M.; Liebman, J . F.; Del Bene, J . E. J . Org. Chem. 1986, 51, 1105. (26) Sletten, E.; Stogard, A. J. Mol. Strurr.: THEOCHEM 1987, 153, 289. (27) Pedersen, C. J. J. Am. Chem. SOC.1967, 89, 7017. (28) Izatt, R. M.; Nelson, D . P.; Rytting, J. H . ; Haymore, B. L.; Christensen, J . J. J. Am. Chem. SOC.1971, 93, 1619.
0022-3654/89/2093-3929$01.50/0 b 1989 American Chemical Societv
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The Journal of Physical Chemistry, Vol. 93, No. 10, 1989
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Computational Details Gradient techniques29 were used to determine geometrical structures of the complexes of azoles with Li+ at the Hartree-Fock level of theory and using the split-valence 3-21G basis set,30whose reasonably good performance for Li-containing compounds is well-do~umented.~~*~~.~',~~ These optimized geometries were then used for single-point calculations at the 6-31G* level33in order to take into account polarization effects. These calculations will be denoted hereafter as 6 - 3 1 G * / / 3 - 2 1 G . The geometries of the bases were taken from ref 13 with the exception of compounds 5, 7, and 9 not included in our previous work. For these three species a similar geometry optimization was performed. The complete set of systems studied in this work is presented in Figure 1. This figure actually shows only those lithiated species which have been found to be stable although, as we shall show later, many other conformations have been in fact studied. Li' binding energies were obtained as the energy differences between the lithiated and the neutral species. The values so obtained are clearly affected by the so-called basis set superposition error (BSSE) which is already significant when dealing with protonation energies34but which is even more important when the attacking ion is Li+. Since we are interested in a comparison between protonation and lithiation energies, we shall pay special attention to the magnitude of this error, which will be evaluated by using the counterpoise procedure of Boys and Bernardi.15 (29) Pulay. P. Applications of Electronic Structure Theory; Schaeffer 111, H. F., Ed.; Plenum: New York, 1977; p 153. Murthaugh, B. A,; Sargent, R. W. H. Compur. J . 1972. 131, 185. Schegel, H. B. J . Comput. Chem. 1982. 3. 214. (30) Binkley, J . S.; Pople. J. A.; Hehre, W. J . J . Am. Chem. SOC.1980. 102. 939. 131) Ikuta. S. Chem. Ph,vs. Lett. 1985, 95. 235. ( 3 2 ) Kaufmann. E.: Schleyer, P. v R. J . Am. Chem. SOC.1985, 107. 5560. (33) Hariharan. P. C.; Pople. J . A . Theor. Chim. Acta 1973, 28, 213. 1 3 4 ) Md. 0.:de Pal. .I. L. G.: Yiriez. M. Theor. Chim. Acta 1988. 73. 307.
Figure 2. Contour m a p of the Laplacian of the charge density for 2Htetrazole protonated a t N4. Positive values of V2p a r e denoted by solid lines and negative values by dashed lines.
To analyze the characteristics of Li+-azole interactions, we shall take advantage of the topological properties of the Laplacian of the electronic density. As it has been shown by Bader,36-38the Laplacian of p is related to the local kinetic G ( r ) and potential V(r) energy densities which appear in the local expression of the virial theorem: ( h 2 / 4 m ) V 2 p ( r )= 2G(r) + V ( r ) (1) Since the kinetic energy density C ( r ) is, by its definition, positive everywhere and V ( r ) is negative everywhere, the sign of the Laplacian of p ( ~ determines ) which of the two contributions dominates in a particular region of space. Accordingly, V2p identifies regions of space wherein the electronic charge of a given system is locally concentrated or depleted. In the first situation V 2 p ( r ) < 0, whereas in the latter V 2 p ( r )> 0. In general, then, negative values of V2p are typical of covalent interactions as found in normal chemical bonds such as N-N, C-H, C-C, etc. Charge is concentrated in the interatomic region leading to an energy lowering associated with the predominance in this region of the potential energy density. On the contrary, interactions between closed-shell systems, as in typical ionic bonds, hydrogen bonds,34 or van der Waals molecules, are characterized by a contraction of the charge toward each interacting system. Consequently, electronic charge is depleted in the interatomic region, leading to a predominance of the kinetic energy density and to a positive value of T2p(r).Therefore, an analysis of the topological properties of V 2 p ( r )will yield direct information on the nature of the interactions between azoles and Li+ ions. The evaluation of V 2 p ( r ) has been implemented by us in the framework of the GAUSSIAN-80 series of programs.40 Finally, it should be mentioned that correlation effects were not taken into account for economic reasons. Nevertheless, we can reasonably assume that, as for other bases,20 inclusion of electron correlation would not significantly change the relative Li+ binding energies reported here, even though these effects are (35) Boys, S. F.; Bernardi, F. Mol. Phys. 1970, 19, 553. (36) Bader, R. F. W.; Essbn, H. J . Chem. Phys. 1984, 80, 1943. (37) Bader, R. F. W.; MacDougall, P. J.; Lau, C. D. H. J . Am. Chem. SOC. 1984, 106, 1594. (38) Wiberg, K. B.: Bader, R. F. W.; Lau, C . D. H. J . A m . Chem. SOC. 1987. 109, 985. (39) Carrol, M.T.; Chang, C.; Bader, R. F. W. Mol. Phys. 1988, 63, 387. (40) Binkley, J. S.; Witheside, R. A,; Krishna, R.; Seeger, R.; De Frees, D. J.; Schlegel, H. 8.; Topiol, S.: Kahn, L. R.; Pople, J . A . Program GAUSSIAN 80; Department of Chemistry, Carnegie-Mellon University, Pittsburgh, PA. The calculation of T 2 p ( r ) was programmed by one of us ( M . A . ) and implemented as a new link of the aforementioned version of the GALSSIAK-80 series of programs. This implementation was appropriately tested by reproducing the results reported in the following articles: Bader, R. F. W.; Slee. T. S.: Cremer. D.: Kraka, E. J . Am. Chem. SOC.1983, 105. 5061. Koch. u' : Frenking. G . : Gauss, J.: Cremer. D.; Sawaryn. A , : Schleber. P. v. R . J . .Ani C h r m S i x 1986. 108. 5732
Molecular Orbital Study of Azole-Li'
Complexes
The Journal of Physical Chemistry, Vol. 93, No. 10, 1989 3931
t c
t Figure 3. Contour map of the Laplacian of the charge density of 2Htetrazole lithiated at N4. Conventions as in Figure 2.
Figure 4. Contour map of the Laplacian of the charge density for 1 Htetrazole-Li+ complex (18). Conventions as in Figure 2.
not zero for absolute Li+ affinities.
is typically greater than 0.5 e- in a protonation process, that involved in the formation of Li+ complexes is always smaller than 0.2 e-. A final significant feature of the azole-Li' complexes is that the N-Li' distance at the equilibrium conformations changes from one system to another, and it is greater the smaller is the corresponding Li' binding energy (see values in the'next section). This is only applicable however to those cases where the complex does not present a bridged structure. When this is the case, we cannot observe a regular trend probably because in these cases, as we shall discuss later, there are other interactions besides the purely electrostatic ones, which play an important role in the relative stability of the complex. Finally, it should be mentioned that when bases with two neighbor nitrogens having a lone pair of electrons are involved, Li+ does not always bridge between both atoms. This finding is in contrast with the behavior observed for other smaller basesZo where the bridged structure is predicted to be systematically the most stable one. Of course, the question of why lithiated pentazole, for instance, does not present a bridged structure or why 4H-1,2,4-triazole only presents bridged structures is still open and will be discussed later on. However, it should be indicated that inclusion of correlation effects might change some of these conclusions, since as found by Del Bene et aL20 for the case of the [F,Li]+ complex, for instance, inclusion of correlation effects leads to a symmetrical bridged form 3.1 kcal/mol more stable than the bent conformation, which is predicted to be the most stable one at the Hartree-Fwk level of accuracy.
Geometries The 3-2 1G optimized geometries for azole-Li' complexes have been summarized in Table I. To facilitate comparison, we have also included those of the neutral az01es.I~ A comparison of the structures of azole-Li+ complexes with those of the corresponding protonated counterparts,13in those cases where the former do not form bridged structures, shows that lithiation effects on the structure of the corresponding azole are similar to those observed upon p r ~ t o n a t i o n but ' ~ much weaker. This is a direct consequence of the different nature of the H+-azole and Li+-azole interactions as revealed by the topological char, are shown in Figures 2 and 3, for the acteristics of V 2 p ( r ) which particular cases of protonated and lithiated 2H-tetrazole chosen as suitable examples. It is evident that while in the protonated species the values of V2p(r)in the internuclear region corresponding to the N-H bond formed upon protonation is negative, showing that, at the equilibrium conformation, the interaction between the incoming H+ and the basic nitrogen is essentially covalent, in the case of the corresponding Li' complex, the value of V2p(r)in the N-Li+ bond region is positive, indicating that electronic charge is depleted in this region and concentrated on both interacting systems: azole and Li'. The latter actually appears as a practically unperturbed spherical charge distribution. This situation is typical of an ionic bond showing that in Li+-azole complexes the interaction is mainly electrostatic and that the N-Li bond is essentially ionic. The topological properties of V 2 p ( r )do not change substantially in those cases where the azole-Li' complex shows a bridged structure as it shows Figure 4 which presents the Laplacian of the density map for complex 18. Again in these cases the interaction is mainly electrostatic and electronic charge is completely depleted in the region between both N lone pairs and Li+; Le., we cannot talk of a Li+ bounded to two lone pairs but simply interacting with both of them. These characteristics of V2p(r)are consistent with th,e fact that while the endocyclic angle centered at the basic center opens considerably upon protonation, it remains practically unchanged upon lithiation. In the first case, as it has been explained elsewhere,*the formation of a new covalent bond between the nitrogen atom and the attacking proton implies a changing from a nitrogen u lone-pair orbital to a u-bonding one; accordingly, the hybridization at the basic center is substantially altered and the endocyclic angle considerably opens. On the contrary, interaction with Li' ions does not lead to the formation of new covalent N-Li bonds and the hybridization of the basic center remains unaffected. Also, the negative charge transfer which takes place at the equilibrium configuration, from the base to the acid, is much smaller when the latter is Li' ion. Actually, while the charge transferred, evaluated by using the Mulliken population analysis,
Li+ Binding Energies We present in Table I1 the total energies of the compounds under study obtained at the 3-21G and 6-31G*//3-21G levels of accuracy. In Table 111 we present the lithiation energies D(B-Li+) (in kcal/mol) involved in the formation of the species indicated, which, as mentioned above, are the only stable ones. As illustrated in Figure 1, only five of them present a bridged conformation while in the remaining complexes Li+ seems to be preferably attached to a single nitrogen lone pair similarly to the corresponding protonated species. It should be noticed that BSSE is quite important at the 3-21G level (about 5 kcal/mol or greater); but. more significantly. these errors are larger (about 7 kcal/mol) for bridged structures. This corroborates previous finding^^'.^^ which show that BSSE increases the stability of lithium-bridging compounds since the AO's of lithium in a bridging conformation can simultaneously supplement (41) Wurrhwein. E. U.; Sen, K . D.: Pople, J . A,: Schleyer, P. v. R. Inorg. Chem. 1983, 22. 496. ( 4 2 ) Kaneti. J.: Schleyer, P. v. R.; Clark, T.; Koa. A. J.; Spitznagel. C . W.; Andrade. J . 6.: Moffat. J . B. J . Am. Chem. SOC. 1986, /OR. 1481.
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57.0 65.8 61.9 45.7 56.8 51.4 49.5 57.1 74.6 58.9 36.7 51.4 44.6
51.8 60.4 56.6 40.9 50.2 44.3 44.4 51.6 66.9 51.9 32.2 45.8 39.3
46.8 54.3 51.0 36.3 44.0 42.0 40.1 45.7 59.9 47.1 30.3 41.9 36.6
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the atomic basis sets of several atoms. Inclusion of polarization functions leads to a decrease of both the absolute binding energies and the BSSE. The first finding is a direct consequence of a more realistic descriptibn of the multipoles of the base and consequently of the ion-base interaction. It should be remarked, however, that although the BSSE at the 6-3 IG*//3-21Ci level is considerably smaller than that calculated at the 3-21G one, it is still significant, being in some particular cases around 2.5 kcal/mol. The consequence is that not only absolute but relative basicities are affected. For example, compounds 11 and 21 present the same lithiation energy at the 3-21G level, but upon correction of the BSSE-much larger in the former because it presents a bridged conformation-complex 11 is predicted to be 1.5 kcal/mol less stable than complex 21. BSSE is also significant regarding the relative stabilities of species 14 and 18 and species 2 and 18. At the higher level of accuracy the error is smaller and its effect on relative stabilities is much less significant than at the 3-21'3 level. In this respect it may be noticed, for instance, that complexes 11 and 21 are predicted to be equally stable before and after correction of the BSSE. Regarding complexes 14 and 18, the situation is similar and the latter is predicted to be always around 1.O kcal/mol more stable than the former. Let us consider now the possible correlation between Li+ binding In the first place it should be taken energies and proton affinitie~.'~ into account that, in principle, the number of lithiated species is greater than that of protonated ones, since while for instance both tautomers of 1,2,4-triazole (12 and 15) lead to a common cation upon protonation that is not longer true when the attacking ion
(E,) and absolute values of the Li' binding energies (D(B-Li') and (b) relative protonation energies (AE,) and relative Li+ binding energies (AD(B-Li')) of azoles. Protonation energies taken from ref 13.
is Li+. Something similar can be said of 1 H- and 2H-tetrazole and of 1H-l,2,3-triazole and 2H-1,2,3-triazole. On the other hand, bridged protonated species are not stable for these systems, so our comparison must be limited to those cases where Li+ is singly coordinated. For these complexes (2, 4, 8, 13, 14, 20, 21, and 23) absolute Li+ binding energies follow the same trend as absolute protonation energies. Actually, we have found that between both sets of values there is a quite good linear relationship (see Figure 5a), with only two exceptions: Li+-pentazole and LP-2H-tetrazole, whose Li+ binding energies are greater than expected from the just-mentioned linear correlation. The existence of this linear correlation is quite significant since it means that covalent interactions in H+-azole systems must be practically constant along the series and that the trend in relative basicities is mainly determined by ion-multipole interactions which are the dominant contributors to the stabilization energy of Li+-azole complexes. Actually, the deviation observed for the case of pentazole and 2H-tetrazole can be easily explained if one takes into account that in both cases the attacked position is directly bonded to another basic center. Under these circumstances a bridged Li+-complex should be expected, and although this is not the case, the Li+-N-N angle is smaller than the corresponding H+-N-N angle in the protonated species.I3 Moreover, if Li' is placed along the line defined by the N-H+ bond of the corresponding protonated azole
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The Journal of Physical Chemistry, Vol. 93, No. 10, 1989
and at a distance of a typical N-Li" bond, the new stabilization energy obtained approaches considerably that predicted from the straight line of Figure 5. More significant is the fact that the Li' binding energies are predicted to be almost 4 times smaller than protonation energies. This result agrees with the experimental findings of Staley and B e a ~ c h a m pwho , ~ ~ found, for a large miscellaneous set of bases, that this attenuation effect is close to 5. In light of our previous discussion it is not surprising to find Li+ binding energies to be smaller than the corresponding proton affinities by approximately a factor of 4. In fact, as we shall discuss a little later, by simple electrostatic effects Li" affinities should be about 3 times smaller that proton affinities. If one takes into account that in protonated species a real covalent bond is formed, while in lithiated ones only electrostatic interactions occur, it is not unreasonable that this factor becomes larger. If one considers relative values, this attenuation factor (see Figure 5b) is around 2; for instance, imidazole is predicted to be 15.5. kcal/mol more basic than pyrazole when the acid is a H + and around 8.5 kcal/mol when the reference acid is a Li' ion at the 3-21G level. More precisely, the slope of the linear correlation in Figure 5b is 1.8 as it should be expected from pure electrostatic interactions. In fact, we can reasonably assume that charge-dipole interactions are dominant and that the polarization term is practically constant along the family of compounds considered. Therefore, taking H" and Li" as point charges, the ratio between relative protonation energies and relative Li" binding energies should vary as
- 2 783yIC \
Y
0 0
Figure 6. Potential energy curve corresponding to the formation of the
Li+-pentazole complex.
R ' ~ / R ~ where R and R' represent the ion-dipole average distance for protonated and lithiated forms, respectively. Taking R and R' to be of the order of the average distance between the attacking ion and the center of the azole ring in the corresponding equilibrium, conformation, Le., around 2.1 and 2.9 A, respectively, which is a reasonable assumption, eq 2 will yield a value of 1.9, in fairly good agreement with the slope of the regression line of Figure 5b. It is also interesting to note that a similar attenuation factor was found14 when studying the NH," binding energies of methyl-substituted pyrazoles and imidazoles. It must be remarked however that this reasoning which is applicable to relative binding energies cannot be generalized to absolute values. In fact, if ion-dipole interactions were the most important contributors to the stability of both kinds of complexes, the ratio between absolute protonation energies and Li' binding energies should be given by the same ratio indicated above. It is then evident that the value actually obtained from our S C F calculations (very close to 4) clearly illustrates that protonated species are stabilized not only by electrostatic effects but polarization, charge-exchange effects, etc., typical of a covalent interaction, are also important contributors.
Stability of Bridged Structures It is reasonable to assume that when bases with two atoms having lone-pair electrons are involved, Li+ may bridge between both atoms, this attachment being symmetrical if the two atoms are equivalent.20 Nevertheless, our results for Li"-azoles show that this is not always the case. We have found that, for instance, when all geometrical parameters of the Li"-pentazole complex are optimized without any restriction, only form 23 is obtained. For a better understanding of this somehow unexpected behavior we have quantitatively obtained the potential energy curve that corresponds to the different paths Li" can follow to approach the azole base. This has been done by fully optimizing the Li+-azole structure for different values of the angles a,& and y defined in Figure 6 (in steps of 10'). We have found that the interaction energy presents a double minimum corresponding to conformation 23, while the bridge structure 24 appears quite high in energy and conformation 25 is about 1.8 kcal/mol less stable than conformation 23. However, in 4H-1.2.4-triazole (15), which presents (43) Staley, R . H.; Beauchamp, J. L. J . Am. Chem. SOC.197597,5920.
- 2 6 2 . 5 53
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Figure 7. Potential energy curve corresponding to the formation of the
Li+-lH-tetrazole complex. the same symmetry as pentazole, only the bridge structure (equivalent somewhat to form 25) is predicted to be stable. We shall discuss later this dissimilarity. At this point it should be said however that the energy gap between conformations 23 and 25 is small enough so that correlation corrections could invert the trend, as it has been foundz0 for [F,Li]+ complexes. Therefore, we can only conclude that at the Hartree-Fock level of accuracy the bridged structure is predicted to be less stable than the bent one, although this conclusion is not affected by inclusion of polarization effects. A similar study was also carried out for 1H-tetrazole. The corresponding energy curve has been plotted in Figure 7. In this case our results show that of the two possible bridged structures only conformation 18 is predicted to be stable, while the other possible bridged conformation, 26, is quite unstable. On the other hand, the complexes lithiated at N2, N3, or N 5 are also considerably unstable with respect to the bridged conformation 18. Since, as we have indicated above, the most important contributions to the stability of Li' complexes are of electrostatic origin, an analysis of the corresponding molecular electrostatic potentials can help in the understanding of these results. Figure 8a shows that the electrostatic potential created by pentazole presents local minima in the neighborhood of each basic center. I t is also evident that the minima close to N2 and N5 are much shallower than those close to N 3 and N4. This would explain why the complex lithiated at N 2 is quite unstable with respect to that lithiated at N3. Similarly, the electrostatic potential of 4H1,2,4-triazole (15) presents two symmetric local minima also close to both basic centers (see Figure 8b). There is however a substantial difference between pentazole and 4H-1,2,4-triazole. In the latter the basic nitrogens are directly bonded to a CH group; their net charges are consequently greater and the minima of the
The Journal of Physical Chemistry, Vol. 93, No. 10, 1989 3935
Molecular Orbital Study of Azole-Li+ Complexes
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Figure 8. Molecular electrostatic potential energy (au) maps of (a) pentazole, (b) 4H-1,2,4-triazole,(c) 4H-1,2,3-triazole, (d) 2H-tetrazole, and (e) 1H-tetrazole. The dashed line shows the position of the Li+ ion in the different conformations of the complex.
corresponding electrostatic potential deeper. The consequence of this is that in pentazole the Li+ ion moves in a region (see Figure 8a) where the two minima are well-separated, Le., moves in a region where the isopotential curves do not connect both minima, and therefore it becomes trapped by one of them yielding form 23. If we consider now 4H-1,2,4-triazole, the two minima are deeper and closer as stated before, and the Li+ ion moves in a region where the isopotential curves cotlnect both minima and therefore the bridged conformation is now expected. An inspection of the molecular electrostatic potentials of 1Htetrazole (17) and 2H-tetrazole (19) (see Figure 8, e and d, respectively) shows that the above discussion is applicable in general. In other words, whenever isopotential energy curves connecting two neighbor minima appear at a distance of the order of the average N-Li+ bond length, the bridged structure will be
formed. On the contrary, if at those distances from the basic center the two minima appear well-separated, the singly bonded structure will be found. It is then obvious that for 2H-tetrazole the latter situation is that found between the two minima and consequently the bridged structure is not stable. For 1H-tetrazole, the bridged structure will only be possible between N1 and N2 but not between N2 and N3, in agreement with our findings. From the same arguments 4H- 1,2,3-triazole (9) should yield both bridged structures (10 and I I ) , as it is in fact predicted by out S C F calculations. Conclusions Our results show that in azole-Li+ complexes the structure of the base changes very little in contrast with the deep changes undergone upon protonation. This is a direct consequence of the different nature of H+-azole and Lic-,i701e interactions '1s
3936
J . Phys. Chem. 1989, 93, 3936-3940 plexes should be the most stable ones, this is not always the case. Pentazole and 1H-tetrazole are significant examples of systems where the expected bridged structures are less stable than the bent ones. From our analysis of the corresponding molecular electrostatic potentials, we can conclude that if the electrostatic minima associated with the two neighboring basic centers are connected by isopotential lines, which are located at a distance from the basic centers similar to or smaller than a typical N-Li' bond length, the bridged structure will be formed. On the contrary, if these isopotential curves appear at distances greater than the average N-Li' bond length, the Li' ion will be trapped in one of the minima yielding a bent conformation.
revealed by the topological characteristics of the Laplacian of the electronic charge density, Li'-azole complexes, if they present either a bent or a bridged conformation, are the result of the interaction of two closed-shell systems, no new covalent bonds are formed, and stabilization arises mainly from electrostatic interactions. One can observe, however, the formation of a new covalent N-H bond in the corresponding protonated species. A consequence of the different nature of H+- and Li+-azole interactions is that the charge transferred from the base to the ion is much smaller in the latter as well as the corresponding stabilization energies. More specifically, Li' binding energies are predicted to be smaller than the corresponding proton affinities by a factor of 4. This implies that there exists a good linear relationship between both sets of values, for those cases where Li' is singly coordinated. On the other hand, the ratio between relative values is about I .8 as it can be expected from a simple electrostatic model, where the dominant term corresponds to ion-dipole interactions. This would indicate that the trend in relative basicities, at least for these kind of bases, is mainly governed by electrostatic interactions. Athough it is reasonable to assume that when bases with two atoms having lone-pair electrons are involved Li' bridging com-
Acknowledgment. This research was partially supported by the DGICYT Project No. PB87-013 1. All calculations were performed at the UAM/IBM and CC/UAM centers, Madrid. Registry No. 1, 288-13-1; 2, 119787-21-2; 3, 288-32-4; 4, 11978722-3; 5, 288-36-8; 6, 119787-27-8; 7, 288-35-7; 8, 119787-32-5; 9, 28834-6; 10, 119787-28-9; 11, 119787-29-0; 12, 288-88-0; 13, 119787-23-4; 14, 119787-24-5; 15, 63598-71-0; 16, 119787-30-3; 17, 288-94-8; 18, 119787-31-4; 19, 100043-29-6; 20, 119787-25-6; 21, 119787-26-7; 22, 289-19-0; 23, 119793-85-0.
Photoselection Studies of CisITrans Isomers of [Ru( bpy),(L),I2+: Interactions in Slnglet Metal-Ligand Charge-Transfer States M. L. Myrick, R. L. Blakley,+ and M. K. De
Evidence for Exciton
Armond*
Department of Chemistry, New Mexico State University, Las Cruces, New Mexico 88003 (Received: August 12, 1988; In Final Form: December 12, 1988)
Steady-state excitation photoselection (SSExP), 77 K absorption, and 77 K emission spectra are presented for cis- and ~rans-[Ru(bpy),(py)~](ClO,), (bpy = 2,2'-bipyridine; py = pyridine), trans-[R~(bpy),(mdpp)~](PF~), (mdpp = methyldiphenylphosphine), fran~-[Ru(bpy)~(tpp)~](CIO~)~ (tpp = triphenylphosphine), and [ R ~ ( b p y ) ~ ( d p p e ) ] ( P(dppe F~)~ = 1,2-bis(diphenylphosphino)ethane) in ethanol glass. SSExP supports the results of the recent interchromophoric coupling model for the SSExP of ruthenium-bipyridine complexes. Independent evidence of excitonic interaction between the chromophores of such species is obtained from the electronic spectroscopy of trans complexes.
Introduction
The steady-state excitation photoselection (SSExP) of [Ru( b ~ y ) ~ (bpy ] ~ ' = 2,2'-bipyridine) and other tris-chelate ruthenium species has been important in the assignment of the lowest excited manifold of this series as spatially isolated localized orbital states, similar to those charge-transfer excited states of monomeric species such as [ R u ( S p y ) ( ~ y ) ~ (py ] ~ += pyridine).' T o date, SSExP results obtained for mono-chelate, tris-chelate, and cis-configuration bis-chelate complexes have been reported's2 with a variety of different solvents, ligands, and counterions. A recent model has been presented that rationalizes the polarization results obtained for the above species,2 based upon the spin-orbit coupling (SOC) interactions of the lowest triplet state with a singlet manifold of mixed (localized/delocalized) character. The results of this analysis compared well with the observed SSExP of bis and tris complexes. In addition, the new model provided a means of relating the SSExP typically seen for most tris chelates to that of the anomalous behavior of [ R ~ ( b i q ) ~ ](biq ~ ' = 2,2'biquinoline).2b This analysis using the interchromophoric coupling (ICC) model depends partly upon the angles between the monomer axis system and dimer/trimer axis systems in complexes that possess multiple chromophores associated with the same core metal ion. If this model is correct, then significant differences are predicted in the 'Present address: Tulane University, New Orleans, LA.
0022-3654/89/2093-3936$01.50/0
polarization properties of bis chelates between cis and trans geometries because of the different angles involved. The purpose of this work is to report SSExP and low-temperature absorption/emission data for some cis/trans isomers and to analyze these data in terms of exciton theory and the ICC model. Experimental Section
The complexes ~ i s - [ R u ( b p y ) ~ ( p y(C104)2, )~] trans-[Ru(bpy)zl (Py)2(C1Od2, frans-[Ru(bpy)2(mdpp)21 (PF6)2 (mdPP = methyldiphenylphosphine), and trans- [ R ~ ( b p y ) ~ ( t p p(ClO4)2 )~] (tpp = triphenylphosphine) were prepared according to literature procedure^.^,^ cis-[R~(bpy)~(mdpp)~]~+ was synthesized but could not be obtained in sufficient purity to permit its use. Luminescence purity of complexes was monitored by the excitation-wavelength independence of the emission. [ R u ( bpy ) ( dpp e) 1 ( P F 6 ) (dppe = 1 ,2-Bis ( d ip h e ny I ph0sphino)ethane). This complex was synthesized by combining 100 mg of [Ru(bpy),Cl,]-3H2O and 400 mg of dppe ligand in 60 ( I ) (a) Carlin, C. M.; De Armond, M. K. Chem. Phys. Len. 1982,89,297. (b) Carlin, C. M.; De Armond, M. K. Chem. Phys. Lett. 1985, 107, 53. (2) (a) Myrick, M. L.; Blakely, R. L.; De Armond, M. K . J . Am. Chem. Sor. 1987, 109, 2841. (b) Myrick, M . L.; Blakely, R. L.; De Armond, M. K. J . A m . Chem. Sor., in press. ( 3 ) Dwyer, F. P.; Goodwin, H . A,; Gyarfa, E. C. Aust. J . Chem. 1963, 16. 544. (4) Durham, B.: Walsh, J. Inorg. Chem. 1982, 21. 329.
0 1989 American Chemical Society
