J . Phys. Chem. 1990, 94, 4796-4804
4796
Experimental and Theoretical Study of Li' Affinities of Methyldiazoles M. Aleam;,+ 0. M6,+M. YBiiez,*>+F. Anvia,%and R . W. Taft*J Departamento de Quimica, C-XIV, Facultad de Ciencias, Universidad AutBnoma de Madrid, Canto Blanco, 28049 Madrid, Spain, and Department of Chemistry, University of California, Irvine, California 9271 7 (Received: October 2, 1989; In Final Form: February I 1990) ~
The gas-phase Li+ affinities of a wide set of methyldiazoles were obtained by ion cyclotron resonance techniques. Simultaneously, Hartree-Fock calculations at the 3-2 1G level have been performed to investigate the structure and stability of these Li+ complexes. A topological analysis of the Laplacian of the electronic charge density reveals that the nature of the N-Li linkage is markedly ionic and therefore noticeably different from the N-H linkage in the corresponding protonated species. However, both experimental and theoretical results show that methyl substituent effects on Li+ binding energies are practically additive. A reasonably good linear relationship between the free energies of adduct formation for Li+ vs H+ is found for this set of compounds, with the slope of this correlation being about 2, Le., similar to that found for unsubstituted azoles. The basicity trends along this family of compounds can be easily rationalized in terms of ion-dipole and polarization interactions. A quantitative analysis of the contributions of these two terms is offered.
Introduction The possibility of measuring equilibrium constants of ionmolecule reactions in the gas phase with high accuracy, by different experimental te~hniques,l-~ has stimulated a growing interest in the study of the gas-phase basicity of a great variety of bases when the reference acid is other than a proton and in particular when the attaching ion is Li+. In this respect, comparison of proton and Li+ affinities reveals4 that, in some cases, relative basicities depend on the reference acid. Furthermore, substituent effects on the strength of a particular base usually depend on the reference acid too. A typical example is provided by methylamines, which present3 a relative ordering of basicity for protonation that is different from that for lithiation. Similarly, the effect of methyl substitution on gas-phase acidities or basicities is not always the same, as it seems to depend upon the nature of the electron-withdrawing groups present in the molecule. For instance, while the substitution of a hydrogen by a methyl group in methanol leads to an increase of its acidityS of about 3.0 kcal/mol, in acetic acid6 this increase is only of about 1 kcal/mol. In our research on the gas-phase basicity of different organic bases we have lately focused our attention on the behavior of azoles and related heterocyclic compound^,^-'^ because they are components of several enzymes and pharmaceuticals and they may be used as model systems to probe the understanding of the coordination chemistry of alkali-metal ions with nucleic acid basesi3 and in the cation-selective transport through biological m e m b r a n e ~ ' ~processes. .~~ Moreover, much of their versatile chemistry arises from their properties as bases. Recently we have studied by theoretical calculations8.I0 the azole-H+ and azole-Li+ complexes as typical polydentate bases. As a further step toward a better understanding of the basicity of these compounds, we have carried out, in the present work, the experimental determination of the Li+ affinities of methylpyrazoles and methylimidazoles, to analyze the possible cumulative effects of the substituents on the basicity of these systems, when the reference acid possesses a Is2 core. This seems a necessary complement to our previous work'* on the H+ affinities of methyldiazoles. In particular we aim to investigate how the basicity increases steadily with increasing methyl substitution and to compare the gas-phase Li+ affinities with the corresponding H+ ones. T h e selected set of mono-, di-, tri-, and tetrasubstituted methylazoles is suited for this purpose, since, on one hand, these compounds contain only one nitrogen as basic center, so bridged structures of the cation which appear for triazoles and tetrazoles are not possible." The basicities of these compounds with respect to either H+ or to Li+ should be sensitive only to the methyl effects. Systematic methyl substitution should result in a smooth increase
'University Universidad Autdnoma de Madrid. of California. t
of the basicity, ideal to perform a systematic analysis of the corresponding intrinsic substituents effects. This particular analysis will be completed by carrying out an a b initio MO study of the Li' complexes between all methylpyrazoles and all methylimidazoles, to determine their structures and to contrast the characteristics of the interactions in Li+-methylazole complexes with respect to those in H+-methylazole systems. In this respect, it is important to establish, without ambiguity, the nature of the bonding between the incoming cation and the basic center in each case. This information should be very useful in the more general interpretation of gas-phase basicities or acidities. Finally, an interesting problem related to the basicity of azoles is the existence, in some derivatives, of tautomeric equilibria. An example is the case of the 3(5)-methylpyrazoles, which are predicted to be equally stable.12 Protonation of these species cannot offer significant information on their intrinsic basicity because they yield the same cation. This is not the case when the attacking ion is Li+, so that an analysis of the calculated Li+ affinities can shed light on the intrinsic basicity of both tautomers. Experimental Section All the chemicals used in this work were available from previous studies, which had been obtained either from commercial sources or were synthesized by previously reported procedure^.'^-'^ The ( I ) Bowers, M. T.; Aue, D. H.; Webb, H. M.; McIver, R. T...Jr. 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, W. A,, 111 Chem. Phys. 1976, 14, 461. ( 5 ) Brauman, J. 1.; Blair, L. K. J . Am. Chem. SOC.1968,90, 5636. Taft, R. W.; Topsom, R. D.; Anvia, F. J . Am. Chem. Soc., in press. (6) Cummings, J. B.; Kebarle, P. Can. J . Chem. 1978, 56, 1. (7) Catalln, J.: de Paz, J. L. G.; YBiiez, M.; Elguero, J. Chem. Scr. 1984, 24, 84. (8) Catalln, J.; Mb, 0.;de Paz, J. L. G.; Perez, P.: Yiiiez, M.; Elguero, J. J . Org. Chem. 1984, 49, 4319. (9) M6, 0.;de Paz, J. L. G.; Ylfiez, M. J. Phys. Chem. 1986, 90, 5597. (IO) M6, 0;;Ylfiez, M.; Elguero, J. J . Org. Chem. 1987, 52, 1713. (11) Alcami, M.; Md, 0.;YlAez, M . J . Phys. Chem. 1989, 93, 3929. (12) (a) Taft, R. W.; Anvia, F.; Taagepera, M.; CatalPn, J.; Elguero, J. J . Am. Chem. SOC.1986, 108, 3237. (b) Catalln, J.; de Paz, J. L. G.; YBAez, 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. ( c ) Catalln, J.; de Paz, J. L. G.; YlRez, M.; Claramunt, R. M.; Ldpez, C.; Elguero, J.; Anvia, F.; Quian, J. H.; Taagepera, M.; Taft, R. W. J . Am. Chem. SOC.1990, 112, 1303. ( I 3) Sletten, E.: Stogard, A. J . Mol. Sfruc. (THEOCHEM) 1987, 153, 289. (14) Pedersen, C. J. J . Am. Chem. SOC.1967, 89, 7017. (15) Izatt, R. M.; Nelson, D. P.; Rytting, J . H.; Haymore, B. L.; Christensen, J. J. J . Am. Chem. SOC.1971, 93, 1619. (16) Elguero, J.; Jaquire, R.; Duc Tien, H. C. N. Bull Chim. Fr. 1966, 3727.
0Q22-3654/9Q/2Q94-4196$02.5Q/O 0 1990 American Chemical Society
Li' Affinities of Methyldiazoles
The Journal of Physical Chemistry, Vol. 94, No. 12, 1990 4797
of the most stable orientations of the methyl groups was carried Fourier-transform ion cyclotron resonance spectrometer employed out. For the sake of conciseness, all results presented and discussed in this work is described elsewhere.20 A I/,-in.-diameter lithium throughout this paper will refer to the most stable conformers. ion source, manufactured by Spectra-Math Inc., was incorporated Li+ binding energies were obtained as the energy differences in the ICR cell. Upon a current of 1.2-1.4 A being passed, between the lithiated and the neutral species. The values so adequate Li+ emission is obtained. Isopropyl chloride is used in obtained are clearly affected by the so called BSSE (basis set small amount ( 1 X lo-' Torr with background pressure of 2 X superposition error), which is already significant when dealing IO-, Torr) to form the monomeric lithium complexes. The with protonation energies30 but which is even more important" metal-ion-transfer equilibria experiments were carried out in this study following those reported earlier by Staley,21+22 B e a ~ c h a m p , ~ , ~ ~when the attaching ion is Li+. Since we are interested in a comparison between protonation and lithiation energies, we shall Gal et a1.,23aand Freiser et Measurements repeated with pay special attention to the magnitude of this error, which will other bases are in good agreement with previous result^.^^^' be evaluated by using the counterpoise procedure of Boys and B e r t ~ a r d i . ~Correlation ~ effects were not taken into account for Computational Details economic reasons. Nevertheless, we can reasonably assume that, Gradient techniques24were used to determine the geometrical as for other bases,26 inclusion of electron correlation would not structures of the complexes of azoles with Li+ at the Hartree-Fock significantly change the relative Li+ binding energies reported here. level of theory using the split-valence 3-21G basis set.2s The The evaluation of zero-point vibrational energies (ZPVE) for reasonably good performance of this basis set for Li-containing systems of this size is economically prohibitive. It would be compounds is well d ~ c u m e n t e d . ~ One ~ - ~ of ~ the problems that interesting, however, to have an estimation of the effect of this arises when optimizing the structures of methyl-substituted azoles correction on calculated Li+ binding energies. For this purpose is that the number of nuclear configurations increases very rapidly we have chosen, as a suitable model system, hydroxylamine bewith substitution. Therefore, to simplify our geometrical model, cause on one hand it is small enough a base to permit the analytical the methyl groups were allowed to adopt only three relative calculation of the second derivatives of the electronic energy with orientations, denoted hereafter as a , b, and c (these orientations respect to the nuclear coordinates, and on the other hand it presents two different basic centers, so it would be possible to know if the ZPVE correction depends on the nature of the basic center. These calculations were carried out with the GAUSSIAN-82 series of programs for the neutral molecule and for the species protonated and lithiated, a t both the oxygen and nitrogen atoms. The results obtained, a t the 3-21G level, show that Z P V E corrections are different for oxygen-protonated and nitrogen-protonated species (7.5 and 9.0 kcal/mol, respectively), but they are much smaller and practically equal (2.0 and 1.8 kcal/mol, respectively) for both U C Li+ complexes. These findings are in agreement with others reported previously in the l i t e r a t ~ r e . ~ ' . ~Therefore ~ we can are schematized for the particular case of N-methylpyrazole-Li" reasonably assume that for Li+-diazole complexes Z P V E corcomplex). In conformations a and b one of the methyl hydrogens rections will not be greater than 2.0 kcal/mol and will not affect lies on the azole plane. In the a arrangement this hydrogen is in a significant way the relative Li' binding energies. Later on toward the Li+ ion, while in b it is away from it. In conformation we shall see that the characteristics of the N-Li+ linkage ratify c none of the methyl hydrogens lie in that plane. Starting from this assumption. these nuclear configurations, all geometrical parameters were The characteristics of Li+-azole interactions were analyzed by allowed to vary and no restrictions were imposed. Even when only the Laplacian of the electronic density. As shown by Bader,33-35 these three relative orientations of the methyl groups are conV 2 p identifies regions of space wherein the electronic charge of sidered, the number of conformers increases noticeably with the a given system is locally concentrated or depleted. In the first degree of substitution (for trimethylated species there will be 27 situation T 2 p ( r )< 0, whereas in the latter V 2 p ( r )> 0. These neutral conformers and 27 lithiated ones). Therefore, for tri- and topological properties of the Laplacian of p are a consequence of tetramethyl-substituted derivatives we have considered only the relationship between this magnitude and the local kinetic G ( r ) orientations a and b, which were found as the most stable in monoand potential V ( r ) energy densities that appear in the local exand dimethyl derivatives. Furthermore, for tri- and tetramethyl pression of the virial theorem: derivatives, the optimization of the azolic ring was carried out, ( h 2 / 4 m ) V 2 p ( r )= 2G(r) V ( r ) exclusively, for the conformer with the methyl groups oriented as in the corresponding most stable monosubstituted derivatives. Since the kinetic energy density C ( r ) is (by definition) positive Then, and without changing the structure of the ring, a search everywhere and V(r) is negative everywhere, the sign of the La-
+
(17) Huball, W.; Pyman, F. L. J . Chem. SOC.1928, 21. (18) Theiling, G. Chem. Ber. 1953, 86, 96. (19) Lions, F.; Ritchie, E. J. J. Proc. R. Soc. N.S. Wales 1941, 74, 365. (20) Catalln, J.; Claramunt, R. M.; Elguero, J.; Mhdez, M.; Laynez, J.; Anvia, F.; Quian, J. H.; Taagepera, M.; Taft, R. W . J . Am. Chem. SOC.1988, 110, 4105. (21) Staley, R. H.; Beauchamp, J. L. J . Am. Chem. SOC.1975, 97,5920. (22) Uppal, J. S.; Staley, R. H. J . Am. Chem. Soc. 1982, 104, 1235. (23) (a) Gal, J.-F.; Taft, R. W.; McIver, R. T., Jr. Spectrosc. Int. J . 1984, 3, 96. (b) Oporti, L.; Tews, E. C.; Freiser, B. S. J . Am. Chem. SOC.1988, 110, 3841. (24) Pulay, P. Applications of Electronic Structure Theory; Schaeffer, H. F., 111, 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. (25) Binkley, J. S.; Pople, J. A,; Hehre, W. J. J . Am. Chem. SOC.1980. 102. 939. (26) Del Bene, J. E.; Frisch, M. J.; Raghavachari, K.; Pople, J. A.; Schleyer, P. v. R. J . Phys. Chem. 1983, 87, 73. (27) Ikuta, S. Chem. Phys. 1984, 108, 441. (28) Ikuta, S. Chem. Phys. Lerr. 1985, 95, 235. (29) Kaufmann, E.; Schleyer, P. v. R. J . Am. Chem. SOC.1985, 107, 5560.
placian of p ( r ) determines which of the two contributions dominates in a particular region of space. In general, then, negative values of C2p are typical of covalent bonds, where 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, positive values of V2p are associated with interactions between closed-shell systems, as in typical ionic bonds, hydrogen bonds,29 or van der Waals molecules, where electronic charge is depleted in the interatomic region, leading to a predominance of the kinetic energy density. Therefore an analysis of the topological properties of V2p(r)will (30) M b , 0.;de Paz, J. L. G.; Ygiiez, M. Theor. Chim. Acta 1988,73,307. (31) Boys, S. F.; Bernardi, F. Mol. Phys. 1970, 19, 553. (32) Meot-Ner (Mautner), M.; Liebman, J. F.; Del Bene, J. E. J . Org. Chem. 1986, 51, 1105. ( 3 3 ) Bader, R. F. W.; Essbn, H.J . Chem. Phys. 1984, 80, 1943. (34) Bader, R. F. W.; MacDougall, P. J.; Lau, C. D. H. J . Am. Chem. SOC. 1984, 106, 1594. (35) Wibern, K. B.: Bader, R. F. W.; Lau, C. D. H . J . Am. Chem. SOC. 1987, 109, 985
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yield direct information on the nature of the interactions between methylazoles and Li+ ions. We have also localized, for the N-Li linkages, the so-called bond critical points, Le., points where the electronic charge density, p, has one positive curvature and two negative curvatures, because the values of p and V2p at these points make it possible to quantitatively characterize the bonding between the azolic nitrogen and the attaching ion. A similar calculation was performed for a selected set of protonated methyldiazoles, for the sake of comparison. The evaluation of the gradient and the Laplacian of the electronic charge density as well as the Hessian matrix has been programmed by one of us (M.A.) and implemented in the framework of the GAUSSIAN-80 series of programs.36
Results and Discussion Structures. For the sake of conciseness we are not going to discuss in detail the optimized structures of the complexes under study, which are available from us upon request. We shall indicate however that as found for azole-Li+ complexes," lithiation effects on the structure of the azolic cycle are so small that one may consider (in good approximation) that the structure of the azole does not change upon lithiation, although it changes appreciably upon protonation. In fact, it is well-known, for instance, that the endocyclic angle centered at the basic center (N2in pyrazoles and N, in imidazoles) opens considerably upon protonation, but it remains practically unchanged upon lithiation. This is so because protonation implies the formation of a true covalent bond between (36) Binkley, J. S.;Witheside, R. A,; Krishna, R.; Seeger, R.; De Frees,
D.J.; Schlegel, H. B.: Topiol, S.;Kahn, L. R.; Pople, J. A. Department of Chemistry, Carnegie Mellon University.
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the basic center and the incoming proton, which results in a considerable alteration of the hybridization of the former and therefore in a significant modification of the azolic ring structure. On the contrary, the interaction between the Li+ and the diazole is basically electrostatic, that is, it is an ionic and not a covalent bond. The topological characteristics of the Laplacian of the charge density corroborate this description. When a methyldiazole becomes protonated, there is a concentration of charge (V2p < 0) in the interatomic region between the basic nitrogen and the proton (see Figure I), indicating that a true covalent bond has been formed. On the other hand, electronic charge is depleted (T2p > 0) between the basic nitrogen and the Li+ ion in the corresponding lithiated species (see Figure 2). This situation is typical of an ionic bond, Le., it is characteristic of a closed-shell interaction. Figure 2 also illustrates that this situation does not change with increasing substitution. Later on we shall analyze, in a little more detail, the characteristics of the V2p maps. Another important feature is related to the relative orientations of the methyl substituents. In this respect it must be taken into account that the overall effect on the stability of the system is not very significant, in the sense that the relative change in energy between one possible orientation of the methyl group and another (about 0.5-0.7 kcal/mol) is very small compared to the total energy. However its effect on relative Li+ binding energies (which are of the order of a few kcal/mol, not greater than 10 kcal/mol) may not be negligible. This effect may be even dramatic when incrementing the number of substituents because the accumulation of these differences may be of the same order or greater than the expected increase in basicity due to the increase in the number of substituents. Consequently, if a systematic analysis of the relative stabilities had not been carried out to determine the most
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4800
The Journal of Physical Chemistry, Vol. 94, No. 12, 1990
stable configurations, the additivity of substituent effects that we found in the theoretical calculations would not have been observed. A further discussion of this point follow toward the end of the next section. Li+ Affinities. W e present in Table I the total energies of methyldiazoles and their Li+ complexes, together with their Li+ binding energies, before and after BSSE correction. Table 11 presents the experimental Lif affinities for the selected set of methyldiazoles considered in this study. The corresponding AGOH+ values40 were also included to facilitate comparison. As indicated above, Z P V E corrections are quite small (-2 kcal/mol) and their effect on relative Li+ binding energies practically negligible. A different matter is the correction due to the BSSE, which is greater in absolute value (-6 kcal/mol) and which is not constant along the series. Actually, it may be observed that there is some kind of saturation effect, since for the tetramethylimidazole, for instance, BSSE is only 0.2-0.3 kcal/mol greater than for dimethyl derivatives. This result implies that BSSE is mainly due to the improvement in the description of the azole when supplemented by the AOs of lithium and that the amelioration in the description of lithium by the atomic basis used to describe the azole does not change significantly when the number of these AOs increases by increasing the number of substituents. I t is important to emphasize that the inclusion of the BSSE does not change the basicity trend along the two families of compounds, but relative to Li+ binding energies become smaller. This is a logical consequence of the fact that, as expected, the BSSE is greater for methyl-substituted derivatives than for the parent compounds. The consequence is that corrected values are in better agreement with experimental measurements, since it is well-known that theoretical calculations, at this level of accuracy, considerably overestimate absolute Li' and H+binding energies and slightly overestimate relative values. In fact, although absolute calculated Li+ affinities are almost twice the experimental ones, the agreement between the relative values is quite good. It should be noticed that this comparison is made between calculated Li+ binding energies and gas-phase free energy changes; however, we do not expect the entropy values to affect in any significant way our conclusions, since changes in the entropy term will be small between different system^.^ Another very important fact is that, as for protonation, both experimental and theoretical Li+ affinities show that methyl substituent effects on the stability of Li+ complexes are practically additive (see Table 111). Besides, no significant saturation effects are found. This behavior may be accounted for by using a quite simple model. As we have illustrated above, the interaction between diazoles and Li+ ions are essentially electrostatic, and the nature of this interaction does not change appreciably with the degree of substitution. To ratify from a quantitative point of view this assertion, we have calculated the value of p and V2pat the N-Li bond critical point for all methyldiazole-Li' complexes studied. The results obtained for methylpyrazoles are presented, as a significant illustration, in Table IV. One may observe that the value of p a t the critical point (i.e.. the maximum value of the electronic charge density between the N and Li nuclei) is quite small and that C2p is positive, independently of the degree of substitution. This explains why ZPVE corrections are much smaller for diazole-Li+ complexes than for diazole-H+ systems and practically constant along the series. Furthermore, these values are almost identical with those reported by Bader and E ~ s i for n ~ CILi, ~ which is a prototype of an ionic bond. The practically constant positive value of V2pa t these critical points along the series indicates that no net charge transfer takes place between the base and the ion, either for the parent compound or for their methyl-substituted derivatives. In other words, one may conclude that in the lithiation process no a-electron delocalization (resonance) effects are manifested and only field and polarizability intrinsic substituent effects play a significant role. This seems to be confirmed by the sharp contrast observed with respect to the N-H bond critical point characteristics for protonated azoles
Alcami et al. TABLE I: Total Energies (au) for Methylimidazoles and Methylpyrazoles and Their Li+ Complexes Li+ binding
substituent H
I-Me 2-Me 4-Me 5-Me I ,2-Me I ,4-Me
1,S-Me 2,4-Me 2,S-Me
neutral Li+ neutral Li+ neutral Li+ neutral Li+ neutral Li+ neutral Li+ neutral Li+ neutral Li+ neutral Li+ neutral Li+
4.5-Me 1,2,4-Me 1,2,5-Me I ,4.5-Me
2,4,5-Me 1,2,4,5-Me
H
I-Me 3-Me 4-Me 5-Me
neutral Li+ neutral Li+ neutral Li+ neutral Li+ neutral Li+ neutral Li+ neutral Li+ neutral Li+ neutral Li+ neutral Li+ neutral Li+
1.3-Me I ,4-Me I S-Me
3.4-Me 3,S-Me
neutral Li+ neutral Li+ neutral Li+ neutral Li+ neutral Li+
4,5-Me l,3,4-Me 1,3,5-Me 1,4.5-Me 3,4,5-Me
neutral Li+ neutral Li+ neutral Li+ neutral Li+ neutral Li+
1,3,4,5-Me
neutral Lit
Imidazoles -223.549 1 1 -230.841 13 -262.363 89 -269.660 29 -262.376 99 -269.671 85 -262.373 80 -269.668 19 -262.374 22 -269.669 57 -301, I90 95 -308.489 59 -301,18853 -308.487 09 -301.188 94 -308.488 14 -301.201 59 -308.498 62 -301.201 91 -308.49961 -301.198 09 -308.495 44 -340.01 5 40 -347.31605 -340.01 5 34 -347.31655 -340.012 77 -347.3 13 79 -340.025 7 1 -347.325 52 -378.83906 -386.141 94 Pyrazoles -223.525 52 -230.803 46 -262.342 52 -269.622 85 -262.351 01 -269.632 16 -262.346 74 -269.626 83 -262.35 1 23 -269.633 65 -301.167 88 -308.45208 -301 .I63 53 -308.446 82 -301.167 33 -308.452 78 -301 .I72 60 -308.455 78 -301. I76 79 -308.462 55 -301.171 72 -308.456 69 -339.989 32 -347.275 29 -339.992 80 -347.28097 -339.987 65 -347.274 56 -339.997 61 -347.284 74 -378.8 13 47 -386.102 90
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64.6
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(see Table IV), where pc is much greater and V 2 p is always negative, indicating a substantial concentration of electronic charge between the incoming proton and the azolic nitrogen, once the equilibrium conformation is reached. Accordingly, one may expect the stability of the corresponding Li' complexes be governed by
Li+ Affinities of Methyldiazoles
T h e Journal of Physical Chemistry, Vol. 94, No. 12, 1990 4801
TABLE II: Gas-Phase Free Energy of Complexation of Lithium to Methyl-Substituted Diazoles (kcalhol) heterocycle std base AGO std base AGo(obs)b 35.4 +o.o 0.1 1,2,4-triazole TH F 36.3 +0.8 + 0.2 pyrazole thiazole +o.o 0.1 36.0 Et20
tetrazoie pyrazole 1-methylpyrazole
4-methy lpyrazole
3( 5)-methylpyrazole 1,4-dimethylpyrazole 1,5-dimethylpyrazoie 3,4,5-trimethylpyrazole
I ,3,5-trimethylpyrazole
1,3,4,5-tetramethylpyrazole
imidazole 1-methylimidazole
2,4,5-trimethylimidazole
MeOCH2CN i-Pr20 Et20 HC02Me HC02Et THF 2-MeTHF octylcyanide acetone octylcyanide acetone n-Pr20 oct ylcyanide acetone octylcyanide 1,3,5-trimethylpyrazole octylcyanide 1,3,4,5-tetramethylpyrazole Me2NCN 4-MePhCoMe Me2NCN I ,S-dimethylpyrazole Me2NCN 3,4,5-trimethylpyrazole Me,NCN c-Pr2C0 4-MePhCOMe DMSO MeCONH2 Me2NCN DMF 1 &naphthyridine I-methylimidazole MeOCH,CH,OH pyridazine 1,2-dimethylimidazole
35.6 38.3 36.0 35.1 36.6 35.4 37.0 40.3
38.0
40.3 38.0 37.5 40.3 38.0 40.3 41.2 40.3 41.8 41.8 40.9 41.8
40.4 41.8 41.5 41.8 41.2 40.9 43.0 42.1 41.8 44.3 46.2 43.0 45.4 44.2 44.6
+ + -0.6 + 0.1 +2.4 + 0.3 +o.o + 0.1 -1.3 +0.4 -1.5
+o.o
+1.9 -0.5 +1.8
0.0 + 0.2
-0.3 +0.4 -1.5 +0.2
+ 0.2
+ 0.2 +0.8 -1.0 + 0.4 +0.1
+0.3 +0.2 +0.6 -0.8
+0.1
+o. 1 +0.8 -0.5
-0.1
+o.o
+ 0.2
-0.8
-0.9 +1.5 +1.1
-1.7
+ 0.2 + 0.2 + 0.4
+0.8
+ 0.5
-0.6 -0.6
+ 0.2
AGOc
35.4 35.5 36.0 36.2 35.9 36.0 36.4 36.3 36.9 37.0 38.4 38.5 38.5 38.0 37.8 39.9 39.5 40.1 40.4 41.3 41.7 41.5 40.7 41.2 41.2 41.7 41.4 41.0 41.7 41.0 43.0 42.9 42.7 42.8 45.1 44.7 44.6 44.8 45.2
av 35.4
AGopd
35.9
205.9
36.0
190.2
36.3
204.7
37.0
208.8
38.5
207.7
38.1
208.0
39.7
212.7
40.3
2 14.0
41.3
216.8
41.2
217.4
41.6
220.5
41.2
215.6
42.8
219.4
44.8
225.1
45.2
225.3
203.0
‘These AGO values for gas-phase stability constants of Li* adducts are taken from: Taft, R. W.; Anvia, F.; Gal, J.-F.; Walsh, S.; Capon, M.; Holmes, M. C.; Hosn, K.; Oloumi, G.; Vasanwala, R.; Yazdani, S. Pure Appl. Chem., in press. bAGo(g)for Li+-transfer equilibrium with indicated standard bases. Positive values indicate greater basicity of the standard AG”(g) for formation of Li* complexes of indjcated heterocyclic base. dAGoH+values taken from ref 40.
ion-dipole and polarization interactions. The first will vary as
R2and the latter as P . In principle, there is an additional term that varies as K3, which is the charge-quadrupole interaction;
however, a suitable choice of the coordinates origin will cancel this term.” It must be taken into account that rigorously speaking the new origin is not unique for the whole series, since it depends3* on the components of the permanent dipole moment. However, since changes in the dipole moment along these series are quite small, one may assume, to a good approximation, that the new origin is the same for the different derivatives and therefore neglect the contributions arising from charge-quadrupole interactions. In good approximation we may consider that the attaching Li+ ion is a point charge and that the polarizability of the diazole is the sum of two components; the polarizability of the azolic ring (which will be constant along the series) plus the polarizability of the methyl groups. With these assumptions the stabilization energies of the Li+ complexes with respect to that of the parent compound (imidazole and pyrazole) will be given by AE(Li+) = 9 / R 2 ( p ’ cos 4’ -
cos 4)
+ 2 a q 2 x ( 1/ r t ) i
(I)
where q is the charge of the attaching ion, which for the case of Li+ complexes can be taken as unity. R is the distance of the ion to the coordinates origin, chosen as indicated above. This distance does not change significantly from one complex to another and,
in very good approximation, it may be taken as constant along the series. p’ is the dipole moment of the methyldiazole considered, and C#J’ the angle between R and the dipole, p and 4 are the corresponding values for the parent compound. LY is the polarizability of the methyl group and ri is the distance between the ion and the i methyl group. Therefore, the summation of the last term of eq 1 runs over the number of methyl substituents (1, 2, 3, or 4 ) . Following our previous arguments, our calculated AE(Li+) values should be correlated by A p = (p’ cos 4’ - p cos 4) and x i l / r 4 . To investigate this point, we have evaluated the A p and the x i l / r 4 terms, for both families. T o simplify the calculation of the latter term, we have used average ri values for methyl groups bonded to N,, C,, and C4,and C5in the case of methylimidazoles and for methyl groups bonded to N,, Cj, C4,and C5in the case of methylpyrazoles. This simplification is easily justified because, in general, the distance between the Li+ ion and the methyl group in monomethyl derivatives does not change appreciably upon further substitution. These average values are summarized in Table V, together with the values of p, p’, 4, and 4’. The corresponding least-squares fitting yield the following equations: for methylimidazoles AE = 5.227Ap
+ 3 1 6 . 6 x ( I / r : ) + 0.10;
r = 0.989 ( 2 )
+ 2 9 8 . 6 x ( l / r t ) - 0.40;
r = 0.981 ( 3 )
i
(37) Kaplan, 1. G. Theory of Molecular Interactions; Elsevier: Amsterdam, 1986. ( 3 8 ) Wangsness, R. K. Electromagnetic Fields; Wiley: New York, 1979; Chapter 8.
for methylpyrazoles AE = 4.817Ap
i
4802
Alcamr et al.
The Journal of Physical Chemistry, Vol. 94, No. 12, I990
TABLE 111: Calculated and Experimental Relative Li+ Binding Energies in kcal/mol AE (es 2) P E ( S C F ) AE,," AEab AEc ACo(exptl) Imidazoles 0.0 0.0 0.1 0.0 0.0 H 0.4 2.6 1.6 2.5 2.1 I-Me -1.2 2.3 1.2 1.4 2-Me -1.8 2.9 1.2 0.9 4-Me 0.3 1.8 1.9 1.4 5-Me 2.6 3.8 3.6 3.6 (3.9)" 1.1 1,2-Me 0.2 3.3 3.6 3.4 (3.4) 1,4-Me 3.3 0.7 4.1 4.2 (4.4) 1 ,5-Me 2.3 (2.3) -2.8 5.1 2.4 2,4-Me 3.1 (3.3) 0.1 2.6 2.8 2,S-Me 3.3 2.9 3.1 (2.8) -0.5 4,S-Me 5.5 4.2 4.4 (4.8) -1.4 1,2,4-Me 3.0 4.9 5.2 (5.8) 1.8 1,2,5-Me 3.6 5.1 4.9 (5.3) 1.4 1,4,5-Me -1.7 5.5 3.9 4.0 4.0 (4.2) 2,4,5-Me 0.0 5.9 6.0 1,2,4,5-Me 5.8 (6.7)
TABLE V Values of p (debyes), Q (degrees), and r (angstroms) Used in EQs 2 and 3 imidazoles pyrazoles IL
3.98 4.41 2-Me 3.83 3.62 4-Me 4.22 5-Me 4.28 1,2-Me 1.4-Me 4.03 4.60 1 ,5-Me 3.49 2,4-Me 4.01 2,5-Me 3.86 4,S-Me 3.93 1,2,4-Me 4.38 1.2,S-Me 4.23 1,4,5-Me 3.67 2,4,5-Me 1,2,4,5-Me 4.01 H I -Me
R 00
Pyrazoles
H I-Me 3-Me 4- Me 5-Me 1,3-Me I ,4-Me I $Me 3,4-Me 3.5-Me 4.5-Me 1,3,4-Me 1,3,5-Me 1,4,5-Me 3,4,5-Me 1,3,4,5-Me
0.0 I .9
1.7 1.3 2.6 3.5 (3.6)' 3.2 (3.2) 4.4 (4.5) 2.9 (3.0) 4.4 (4.3) 4.2 (3.9) 4.6 (4.9) 5.9 (6.2) 5.3 ( 5 . 8 ) 5.2 (5.6) 6.7 (7.5)
0.0
0.0
-0.6 -I .3 0.4 2.2
2.4 2.7 0.3 0.4
-1.8
5.1
-0.2 1.7 -0.9
2.8 2.8 3.0 3.0 0.7 5.4 5.4 3.1 3.4 5.8
1.0
2.5 -1.4 0.5 1.8 0.9 0.5
0.4 2.2
ri
5.39 3.45 3.22 5.44
H I-Me 3-Me 4-Me 5-Me 1,3-Me 1,4-Me 1,5-Me 3,4-Me 35Me 4,5-Me 1,3,4-Me 1,3,5-Me 1,4,5-Me 3,4,5-Me 1,3,4,5-Me
P
4
2.50 2.58 2.18 2.51 2.98 2.23 2.53 3.02 2.23 2.67 2.98 2.21 2.68 2.94 2.65 2.3
14.0 26.9 9.8 7.5 14.0 22.8 19.5 23.3 2.7 10.5 9.0 16.7 20.3 18.4 6.1 15.9
Ti
3.33 3.25 5.45 5.42
,
0.0
0.7
1.8
1.d
1.1
2.2 1.8'
3.0 3.7 3.0 4.9 2.5 4.4 3.6 4.4 6.3 5.3 4.7 6.7
4
6.4 8.9 13.3 4.6 1.7 13.1 7.6 4.5 12.2 8.6 0.5 19.8 9.9 4.8 7.1 9.5
3.4 4.0
4.9 5.O 5.3
OAE, corresponds to the first term of the right-hand side of eqs 2 and 3. " A E , corresponds to the second term of the right-hand side of eqs 2 and 3 . cFor imidazoles, A E = AE, + AE, + 0.1; for pyrazoles, AE = AE, + AE, + 0.4. Values in parentheses were predicted assuming constant increments of 2.5, 1.4, 0.9, and 1.9 kcal/mol for I-, 2-, 4-, and 5-methyl substitution, respectively. CValues in parentheses were predicted assuming constant increments of 1.9, 1.7, 1.3, and 2.6 kcal/mol for I-, 3-, 4-, and 5-methyl substitution, respectively. /Experimentally is not possible to distinguish between 3-methyl- and
5-meth ylpyrazole.
TABLE IV: Values of the Electronic Density ( Aand the Laplacian of the Electronic Density (V*p,) at the N-Lf+ Bond Critical Points of Methylpyrazole-Li+ Complexes (All Values in au) Li+ H+ Pc V2PC Pd V2P,' H 0.0413 0.2961 0.3006 0.3190 -1.5056 0.0422 1-Me 0.2994 0.3181 -1.5133 0.0421 3-Me 0.2985 0.3177 -1.5178 0.0418 4-Me 0.3017 0.3182 -1.5043 0.0422 5-Me 0.0427 0.3021 1.3-Me 0.0426 1,4-Me 0.3027 0.0429 0.3049 1 ,5-Me 0.301 7 3.4-Me 0.0425 3,S-Me 0.0429 0.3040 0.0427 0.3049 43Me 1,3,4-Me 0.0429 0.3032 1,3,5-Me 0.0432 0.3049 1,4,5-Me 0.0433 0.3073 3.4.5-Me 0.0430 0.3050 0.3078 1,3,4,5-Me 0.0436 a Values corresponding to the N-H+ bond critical point of the protonated species.
The goodness of these multivariate correlations is put in evidence in Figure 3, where we have plotted the AE values predicted from
0
E 6 00
44
00
a a L.
-3
2 00
w
a 0 00
0 00
2 00
4 00
-
7-77
T
6 00
1
-
7
-
m
8 00
aE(l.i+) SCF (kcal/nrol) Figure 3. Relative Li+ binding energies predicted by using eqs 2 and 3 vs SCF calculated values for (A) methylimidazoles and ( 0 )methylpyrazoles. eqs 2 and 3 vs the SCF values (see Table 111). We conclude that the ion-dipole and polarization interactions are the basic contributors to the stabilization of the Li+-methyldiazole complexes. There are however some additional facts that confirm this conclusion. Following our previous reasoning, the coefficient of the first term of eqs 2 and 3 should be equal t o q / R 2and that of the second term should be proportional t o the polarizability of the methyl group. A quite simple calculation, setting q = 1, yields a value for R 3.5 A, which is not only reasonable but also consistent with the structure of the complex. T h e polarizability estimated for the methyl group is 1.81 A3, which is in remarkably good agreement with the experimental estimated value (1.76 A)) one can obtain by subtracting the polarizability of ammonia from that of m e t h ~ l a m i n e . ~ ~ Another important feature is that both contributions are of the same order of magnitude. This implies that the additivity of methyl substituent effects is not only a consequence of the expected additivity of the polarizabilities of the methyl groups. Changes induced by the methyl substitution on the direction and magnitude of the diazole dipole moment yield ion-dipole interaction energies that are also additive. Furthermore, a perusal of Table I11 shows that in several cases the ion-dipole term clearly dominates and that the contributions of the two terms are, in some cases, opposite in sign. Obviously the polarization term is always stabilizing, but this is not the case for the ion-dipole term whose sign depends on the relative orientation of the dipole. (39) Hirschfelder, J . D.; Curtis, C. F.; Bird, R. B. Molecular Theory of Gases and Liquids; Wiley: New York. 1954. Applequist, J. P.; Cook, J . R.; Fung, K.K.J . Am. Chem. SOC.1972,94,2952. Zeiss, G. D.; Meath, W.J . Mol. Phys. 1977, 33, 1155.
The Journal of Physical Chemistry, Vol. 94, No. 12, 1990 4803
Li+ Affinities of Methyldiazoles
3
1 . 1.2.4 triazole 2 yrazole 3 . Phiazole 4. 4-Me yrazole 5. 3(5>-& ovrazole 6. 1-ke p&&ole 7 . 1 . 4 - d i - k pyrazole 8. 1.5-di M e pyrazole
1 1 . 1,3.5-tri Ye pyrazole 1 2 . 1 -Ne imidazole 13. 1,3,4,5-tetra Me p 14 1,2-di Me imidazoxazole 15. 2.4.5-tri M e imidazole
9
imidazole 10 3.4.5-tri Me pyraz ole
~
188.0 30.0
36.0
AGO ( L i + )
42.0
48 0
(kcal/mol)
Figure 4. Plot of AC"(H+) vs AC"(Li+) for methyldiazoles. ACo(H+) corresponds to the process azole(g) to azole(g) Li+(g) ;ri azole-Li+(g).
+
This partitioning is particularly relevant for those systems involved in tautomeric equilibria. Let us consider, as a suitable example, the case of 3(5)-methyl pyrazoles. Both present the same proton affinity since they are equally stable and yield the same cation upon protonation. However, when the reference acid is Li+, 5-methylpyrazole is predicted to be 1 kcal/mol more basic than 3-methylpyrazole. But the most important fact is that their increased basicities with respect to the parent compound have quite different origins. Our results show that the ion-dipole interactions in 3-methylpyrazole would yield a decrease of its basicity with respect to pyrazole. This decrease is surpassed by a greater stabilizing effect due to the polarization of the methyl group, which in this case is close to the Li+ ion. On the contrary, the increase of the basicity of the 5-methylpyrazole with respect to the parent compound is mainly due to a favorable ion-dipole interaction, while the stabilization due to the polarization term is much weaker than in 3-methylpyrazole, because the methyl group is further away from the Li+ ion. In general we have found that when the methyl group is a to the basic center, the ion-dipole term is negative and the increase in basicity is essentially due to the polarization term. Quite on the contrary, when the methyl group is further from the basic center, the increase of basicity is basically due to the ion-dipole term, which is always positive. Correlation between LI+ and H+ Affinities. Let us now consider the possible correlation between Li+ and proton affinities. In Figure 4 we have plotted the AGa(H+) vs ACo(Li+) for the 13 compounds of this series considered in this study, together with 1,2,4-triazole, tetrazole, and thiazole.40 This figure clearly illustrates that the correlation is quite good, with the exception of tetrazole. The nearly 6 kcal/mol greater stability of the Li+ complex than is predicted by this correlation points to the possible formation of a bridged Li+ structure which has been predicted" as the most stable conformation of the tetrazole-Li+ complex. The mechanism by which this bridged structure is formed is under investigation. The values of AGO in Figure 4 are uncorrected for relatively small rotational TAS effects, which are known3 to be within f l kcal/mol. These nonelectronic effects, which are too small to camouflage major electronic effects, will be the subject of a subsequent publication. The goodness of the AG"(H+) vs ACo(Li+) correlation for methyldiazoles indicates that methyl (40) Gas-phase basicities (AGH+ values) and proton affinities for these series are summarized in ref 1 I C and for 1.2.4-triazoleand tetrazole in: Taft. R . W.; Anvia, F.; Catalan, J.; de Paz, J. L.G.; Elguero, J., to be published.
+ H+(g) 9 azole-H+(g)
and ACo(Li+)
substituent effects have much the same origin for both kind of processes. It is also obvious that absolute proton affinities are about 5 times greater than Li+ affinities, in agreement with previous findings reported in the literature3 for a large miscellaneous set of bases. This is not surprising in the light of our previous discussion, in the sense that in protonation process a new covalent bond is formed, while lithiation implies only a closed-shell interaction. More interesting is the fact that the slope of the linear relationship of Figure 4 is around 2.4. It is striking that this slope is very similar to that reported for azoles,l' where only the iondipole interactions should be important contributors to the relative stabilization of the complex, since the polarizability along the series may be considered practically constant. In that case, it has been shown" that this slope was approximately given by [R(Li+),/R(H+)*] E 2.0, where R(Li+) and R(H+) are the distances from the Li+ and the H+ to the dipole, respectively. On the other hand, the good linear relationship of Figure 4 indicates that intrinsic resonance effects, present only in protonation, are either constant or steadily increase with methyl substitution. In the first case, the inclusion of this term will not change the slope of the linear correlation of Figure 4, and for our purpose we may consider that relative protonation energies should fulfill an equation similar to (1):
AE(H+) = (q/R(H+)')(p'
COS
4' - p
COS
4)
+
2 a q 2 C (1 /ri(H)4)
+ constant
(4)
i
where the only difference is that it includes a constant term that accounts for the resonant effects and that, as indicated above, R ( H + ) is now the distance from the incoming proton to the coordinates origin and r,(H) the distances from this proton to the methyl substituents. According to this, the ratio between relative protonation energies and relative lithiation energies should be given by AE(H+)/AE(Li+) = R(Li+)'/R(H+), - [2aq2/AE(Li+)] X tC(1/ri(H)4) - ( R ( L i + ) 2 / R ( H + ) 2 ) c ( 1/ri(Li)4)1 (5) i
I
Since the N-H+ distances in protonated diazoles are considerably shorter (about 0.9 %(. shorter) than N-Li+ ones in complexes Li+-diazole, R ( H + ) and r,(H) should also be shorter than those for lithiated species (see Table V). A trivial calculation shows then that the expression within brackets in eq 5 is practically zero (the l / r ( H ) 4 term is about twice the 1 r(Li)4 one, but the latter is multiplied by the ratio R(Li),/R(H) , whlch is also about 2.0).
1 '
J . Phys. Chem. 1990. 94, 4804-4809
4804
The consequence is that eq 5 reduces, in good approximation, to (6) AE(H')/AE(Li+) N R(Li+)2/R(H')2 N 2.0 explaining why the ratio between relative protonation energies and relative lithiation energies for azoles and methyldiazoles are practically equal, even though in the latter, polarization contributions are important while in the former they are practically negligible. Equation 6 also shows that the slope predicted, under the assumption that resonance effects are constant along the series, is slightly smaller that the experimental one, indicating that these effects, in fact, increase with methyl substitution. In summary, the difference between relative protonation affinities and relative Lie affinities are due not only to stronger ion-dipole and polarization interactions in the former but also to the presence of methyl T-electron donor resonance effects that do not take place in the latter.
Conclusions W e have found a good linear relationship between H + and Li+ affinities of methyldiazoles, although the latter are more than S times smaller than the former. This is a consequence of the weak interaction base-acid when the reference acid is Li'. A detailed analysis of the topological characteristics of the Y2p for the
diaLole-Li' complexes shows that this is a closed-shell interaction, dissimilarlq to that in protonated species, which leads to the formation of a covalent bond between the azolic nitrogen and the incoming proton. Furthermore, the value of V2p at the N-Li bond critical point does not change significantly upon substitution, so one may conclude that only electrostatic and inductive interactions are responsible for the concomitant increase in the gas-phase basicity. In fact a simple model including only ion-dipole and polarization terms accounts for the observed substituent effects on L.i+ affinities. This model also shows that both terms contribute to this increase and that the polarization contribution is not always the dominant one. 5-Methyl- and 4,5-dimethylpyrazoles and 5-methyl and 1 ,S-dimethylimidazoles are clear examples of this situation. In summary, absolute and relative proton affinities of methyldiazoles are greater than the corresponding Li' affinities not only because in the former case ion-dipole and polarization interactions are stronger but also because in the latter there are no resonance effect contributions to the stabilization of the complex.
Acknowledgment. This research has been partially supported by the DGICYT Project No. PB87-0131. W e thank a referee for helpful suggestions.
Chromophoric Fine Tuning and the Interchromophoric Coupling Model in Ruthenium(I I ) Polypyridyl Complexes R. L. Blakley,+M. L. Myrick, R . Pittman, and M. K. De Armond* Chemistry Department, New Mexico State Unicersity, Las Cruces, New Mexico 88003 (Receicrd: October 9, 1989)
The interchromophoric coupling model ( J . Am. Chem. SOC.1988, 110, 1325) succeeded in rationalizing qualitatively and quantitatively the excited-state absorption and emission properties for true mono-, bis-, and tris(diimine) chelates of Ru(I1). Specifically, the weak interligand coupling between identical ligands for bis and tris complexes results in regions of localized and delocalized excitation for the singlet MLCT excited state. Spin-orbit coupling then produces localized orbital emitting states of mixed polarization. The magnitude of the P,,, (maximum excitation photoselection) is typically 0.23 for true tris and 0.34 for true bis compounds and approaches 0.5 for mono(diimine) complexes. I n every case, the magnitude of P,,, depends upon the number of emitting degenerate states. These new results present data for mixed-ligand chelates in which the emitting state's degeneracy is systematically reduced from that of the true parent bis complexes. This systematic reduction in the interligand coupling between the near degenerate levels of the diimine complexes results in P,,, values ranging from 0.20 to 0.32 in a continuous fashion and consistent with the predictions of the ICC model.
Introduction Many models for the emitting state of ruthenium polypyridyl complexes have been proposed. Early work by Crosby et al.' resulted in a three-state model based upon an excited state of D3 symmetry-the full symmetry of the ground state of [ R ~ ( b p y ) ~ ] ~ ' (bpy = 2,2'-bipyridine). This analysis was based upon mathematical interpretation of luminescence lifetime and quantum yield measurements as a function of temperature and assumption of fuil thermal equilibration of the emitting levels of the complexes down to approximately 1.4 K. the lowest temperatures attained. More recent data, however, have indicated that a t least the lowest excited state of the molecule must possess symmetry lower than D,. These data have been obtained from p h o t o ~ e l e c t i o n ,ex~~~ cited-state resonance R a m a r ~ , ~excited-state -~ ab~orption.~-lexcited-state circular dichroism,I2 and excited-state ESRIj methods and appear to indicate that a description of the lowest state in terms of C,, or Cz symmetry is more appropriate. Attempts to make a multilevel assignment with a C, description of the excited state as originally proposed (necessitating three 'Work done at North Carolina State University. *To whom correspondence should be addressed. 0022-3654/90/2094-4804%0250/0
emitting levels a t 0, IO, and 60 cm-I) have met with little success. Kober and MeyerI4 reinterpreted the original data of Crosby in ( I ) dosby, G. Ace. Chem. Res. 1975, 8, 231. (2) (a) Carlin, C.; De Armond, M. Chem. Phys. Lerr. 1982,89, 297. (b) Carlin, C.; De Armond, M. Chem. Phys. Leu. 1985, 107, 53. (3) (a) Myrick, M.; Blakely, R.; De Armond, M.; Arthur, M. J . Am. Chem. Soc. 1988,110, 1325. (b) Myrick, M.; Blakely, R.; De Armond, M . Chem. Phys. Lett., in press. ( 4 ) Dallinger, R.; Woodruff, W. J . Am. Chem. Soc. 1979, 101, 4391. ( 5 ) Carroll, P.; Brus, L. J . Am. Chem. Soc. 1987, 109, 7613. (6) McClanahan, S.: Dallinger, R.; Holler, F.; Kincaid. J . J . A m . Chem. Soc. 1985, 107, 4853. ( 7 ) Kumar, C.; Barton, J.; Gould, I.; Turro, N.; Van Houten, J . Inorg. C-hem. 1988, 27, 648. (8) (a) Anderson, D.; Orman, L.; Chang, Y.; Yak, T.; Hopkins, J. J . Am. Chem. Sur.. in press. (b) Orman, L.: Hopkins, J . Chem. Phys. Lerr. 1988, 1 4 9 , 375.
( 9 ) Braterman, P.; Harriman, A.; Heath, G.; Yellowlees, L. J . Chem. Soc., Dalton Trans. 1983, 180 I . (10) Milder, S.; Gold, J.; Kliger, D. J . Phys. Chem. 1986, 90, 548. ( 1 1 ) Hauser, A.; Krausz, E. Chem. Phys. Lerr. 1987, 138. 355. (12) (a) Gold, J.; Milder, S.; Lewis.; Kliger, D. J . Am. Chem. Sot. 1985, 107, 8285. (b) Milder, S.; Gold, J.; Kliger, D. J . Am. Chem. Soc., in press. ( 1 3) Yamauchi, S.; Komada, Y.: Hirota, N. Chem. Phys. Letr. 1986, 129, 197. (14)
Kober. E.; Meyer, T. fnurg. Chem. 1984, 23,
C 1990 American Chemical Society
3877
