J. Phys. Chem. 1903, 87, 1833-1834
I I
40
I I I I
50
I
Flgwe 4. Cesium133 chemical shift vs. (18-crown-0)/(CsBPh,) mole ratio and temperature in ammonia; (CsBPh,) = 0.001 M. Solid lines are calculated curves. Dashed lines connect points at mole ratios of 0 and 1. Inset: Ceslum-133 chemical shift vs. concentration of CsBPh, in liquid ammonia at 6.0 OC.
interact with the cesium cation in the 2:l complex. Therefore, a linear temperature dependence was also assumed for 8cscz+.x-according to
(6cBcz+.x-)t = ( ~ c ~ c Z + . ~ - ) 2+5 "b(t c - 25OC)
(6)
Four parameters, Kx2 AHox2, (8cSc+ . x - ) ~ ~and ~ c ,b were adjusted. The results are given in $able I. The average standard deviation of the chemical shift (a, = 0.42 ppm) is about twice the experimental error. The mole ratio dependence of the chemical shift of CsBPhl in ammonia was also studied a t a total concentration of 0.0075 M of the salt a t 14.5 "C. The value of
1033
Kx2 obtained from these results was less than half the corresponding value given in Table I for (Cs+) = 0.001 M at the same temperature. These results indicate that ion association is important in liquid ammonia and would have to be included in a complete treatment. The "thermodynamic parameters" obtained for ammonia solutions are, therefore, conditional in nature. It appears that the formation constant for the 2:l complex in liquid ammonia is much larger than in methylamine. If ion association of the 1:l and 2:l complexes could be ignored, then one would expect the reverse order since ammonia is a better donor solvent than methylamine. The limiting chemical shift of the 2:l complex at high values of R in methylamine is almost temperature independent and practically equal to the values in other nonaqueous solvent^.^ Surprisingly, the limiting chemical shift in ammonia is more than 80 ppm downfield as compared to other solvents. This large downfield chemical shift could be due to the interaction of the cesium cation in the sandwhich complex with the solvent, since the chemical shift of the free cation in liquid ammonia is larger than 122 ppm. A more complete description of the formation of the 2:l complex in ammonia would require determination of the ion-association parameters of both the salt and the 1:l complex. Formation of the 2:l complex (equilibrium 3c) is enthalpy stabilized but entropy destabilized in both solvents. The entropy of formation of the 2:l complex seems to be nearly anion independent, but strongly solvent dependent. However, the enthalpy of the complexation reaction is both anion and solvent dependent. In methylamine solutions, differences in the stability of the 2:l complexes of cesium iodide and cesium tetraphenylborate with 18-crown-6 are mainly determined by the enthalpy contribution to the free energy of formation. The larger complexation constant for cesium tetraphenylborate compared to that for cesium iodide reflects the difference in the degree of ion association of their corresponding salts and complexes. The much more positive entropy of formation in liquid ammonia might be due to stronger solvation of the cesium cations by ammonia molecules than by methylamine molecules.
Acknowledgment. We gratefully acknowledge the support of this work by National Science Foundation Grants CHE-80-10808 (A.I.P.) and DMR 79-21979 (J.L.D.) Registry No. CsI, 7789-17-5; CsBPh,, 3087-82-9.
COMMENTS Decarboxylationof Acetate Ions
Sir: Electron spin resonance spectroscopic studies have revealed many details about radiation processes in recent year~.l-~In particular, the use of single crystals has shown (1) H. C. Box, 'Radiation Effects E.S.R. and ENDOR Analysis", Academic Press, New York, 1977. (2) M. C. R. Symons, Pure Appl. Chem., 53, 223 (1981). (3) S. Ya. Pshezhetakii, A. G. Kotov, V. K. Milinchuk, V. A. Rozinskii, and V. I. Tupikov, 'E.P.R. of Free Radicals in Radiation Chemistry", Wiley, New York, 1974. 0022-365418312087-1833$01.50/0
that neutral or ionic fragmenta are often trapped in specific sites with specific orientations, a general finding being that changes relative to the original geometry are often small. When a primary radical product undergoes unimolecular decomposition, the resulting fragments are usually expected to remain in their original sites. A clear example of this is the detection of intermolecular hyperfine coupling between the radical and nonradical fragments, as found, for example, for alkyl radical-halide ion adductsS4+ (4) E. D. Sprague and F. Williams, J . Chem. Phys., 54, 5425 (1971).
0 1983 American Chemical Society
1834
The Journal of Physical Chemistry, Vol. 87, No. 10, 1983
While ESR and ENDOR spectroscopy can often provide useful information, it seems that the electron spin-echo technique can give even more detailed geometric information. Kevan and his co-workers have exploited the power of this method to great effect in their studies of the local geometry of the matrix surrounding radical ions, and in particular their work has provided conclusive evidence for the vacancy model for trapped Kevan and his co-workers have recently probed the environment of methyl radicals formed in the radiolysis of lithium acetate dihydrate at 77 K using this techniquelo and have drawn attention to their apparently surprising result that these radicals must have moved inward, toward the C02site rather than in the opposite direction as might have been expected for a fragmentation process. A previous ENDOR study by Iwasaki and co-workers" had established that any movement away from the original site of the methyl group must have been small, but they were not able to be very specific. The spin-echo probe was able to show that the six near-neighbor lithium ions are divided such that two interact with the electron slightly more strongly than the remaining four. This has to mean a movement of ca. 0.03 nm towards the original site of the COS group. The aim of this note is to suggest that this is actually the expected result for decarboxylation of the acetate radical. This is thought to be important in that it underlines the accuracy of the spin-echo method, in addition to providing detailed mechanistic insight. The normal primary centers in irradiated carboxylates are RC02. and RC022-. It has long been realized that RC02. radicals under a unimolecular dissociation RC02.
-
R.
+ C02
(1)
whereas RC02- anions are not normally a source of Rradicals. The SOMO for R-C02. radicals comprises a combination of 2p oxygen orbitals. There are three lowlying electronic states available, the one that correlates with bond breaking being the symmetric 2A' state for which the LOCO angle opens, and there is only a very shallow minimum for the C-C potential we11.12 As the bond breaks, so the C02fragment becomes linear and the CH, fragment becomes planar, as indicated in Figure 1. It is the high stability of linear C02 that provides the driving force for this reaction. I suggest that there will be little tendency for the oxygen ( 5 ) S.P. Mishra and M. C. R. Symons, J . Chem. SOC., Perkin Trans. 2, 391 (1973). (6) M. C. R. Symons and I. G. Smith, J . Chem. SOC.,Perkin Trans. 2, 1362 (1979). (7) D.-P. Lin and L. Kevan, Radiat. Phys. Chem., 17, 71 (1981). (8) M. Narayana and L. Kevan, J . Am. Chem. SOC., 103,1618 (1981). Feng and L. Kevan, Chem. Reu., 80, 1 (1980). (9) D.-F. (10) M. Narayana, L. Kevan, and S.Schlick, J. Phys. Chen., 86,196 (1982). (11)T. Toriyama, K. Nunome, and M. Iwasaki, J . Chem. Phys., 64, 2020 (1976). (12) S. D. Peyerimhoff, P. S.Skell, D. D. May, and R. J. Buenker, J . Am. Chem. SOC.,104, 4515 (1982).
Comments
'
H---*
I /o
1
I
i HH
0,
Figwe 1. Suggested relaxation for the CH, and COPfragments when CH,-CO,* radicals break down in lithium acetate. The distances, given in angstroms, are based on the assumption that the oxygen atoms remain fixed during the relaxation.
atoms to move from their original sites. In the acetate ions these oxygen atoms interact strongly with Li' neighbors and will be very precisely located. Although the negative charge has been greatly reduced, there will still be little freedom of motion for these atoms. Thus, the central carbon will move along the b axis as shown in Figure 1, away from the methyl radical. In contrast, on flattening, it will be the small, light protons of -CH3that will move. Hence, this radical finds itself as constrained on one side to about the same extent as in the parent anion, but with a small void on the other side provided by the shifted carbon atom of the C 0 2 molecule. It seems clear that there should be a very minor relaxation of methyl in that direction, namely, toward C02 rather than away from it, as observed. A simple calculation suggests that the carbon of C 0 2 must move ca. 0.064 nm. Thus, a movement of 0.03 nm in the same direction for the methyl radical is eminently reasonable. It is of interest to compare this result with our recent resultd3for RN02+.radical cations, which are isoelectronic with RC02*radicals. The reaction analogous to reaction 1 would give R. + NO2+. However, there was no trace of R. radicals. In addition to low concentrations of the parent cations, we obtained high yields of species identified as (ONOR)+radicals by ESR spectroscopy. The analogous reaction, giving OCOR radicals, does not seem to have been detected in studies on RC02. radicals, the only detected process being reaction 1. Acknowledgment. I thank Professor L. Kevan for encouraging comments. Department of Chemistry The University of Leicester Leicester, LE1 7RH England
Martyn C. R. Symons
Received: December 28, 1982
(13)D.N. R. Rao and M. C. R. Symons, Tetrahedron Lett., in press.
