282
NOTES
formation, which is (in most cases) in equilibrium with the subsequent slower step or steps. Values of the equilibrium quotient for ion pair formation, Kan (Ltf-l), depend mainly upon charge type. In normal, llonhindered complexation the slow step in complex formation is the first elimination of a coordinated water molecule, described by the rate constant kon (sec-l) for the nth step. For normal complexation the observed rate constant is lelated to the mechanistic constants according to kn SnbzKan (2) where S, is a statistical factor depending upon the number of available bonding sites in the coordination shell of the metal ion. When n = 1, kon is approximately (1-3) X lo4 sec-’ for nickel(I1) and (3-10) X IO6 sec-l for cobalt(I1). This value is essentially independent of ligand proper tie^.^ For the fully aquated dipositive metal ion reacting with a singly negative 2 M - I and = 1. From eq 2 the comanion, K,I puted values of kol, using the results in Table I1 and the above constants, are: nickel(II), kOl = 1.3 X lo4 sec-’; cobalt(II), kol < 5 X 106 sec-l, These figurcs are consistent with the values cited.’ The situation with regard to higher order complexation is less well understood than the corresponding case for reaction of fully aquated metal ions. Nevertheless, it is here that the picolinic acid results are most interesting. We calculate using Xz = 2/3, Sa = Ka2 = 1, and Kaa = 0.3 the following values of ko2 and k03 (sec-l): nicld(I1): kO2= 1.8 x 106, koa = 1.6 X lo6; cobalt(I1): k02 = 3.6 X lo7, koa = 1.9 X lo7. It is therefore concluded that replacement of two or more inner coordination sphere waters by picolinate enhances the rate of water loss, This effect may arise, in part, from the negatively charged carboxylate ligand. That is, bonding to -COO-should reduce the positive charge density on the central metal ion, thereby loosening the binding to the remaining waters. Studies of a-amino acids have yielded similar results.4r8 Evidence is also available, however, that other factors are involved. Consider, for example, the ligand pyridine-2,6-dicarboxylate (pydic), For the reaction Nipydic pydic2-, k2 = 2.4 x lo4 M-l sec-1.6 With XZ = and K Z = 0.3, we obtain ~ O = Z 1.6 X lo6 sec-’ for this ligand, almost equal to the picolinate k02, despite the presence of an additional -COO- group. Clearly, the formal charge of the bound ligand cannot be a most significant factor in these systems in determining the rate of water loss from the inner coordination sphere of the complexed metal ion, Yet the bound picolinate anion i$ plrying an important role in determining the rate constant for this process, since ~ O Z / ~ O I= 12. Quite likely the enhancement is a result of the interaction of delocalized 7r orbitals on the ligand with T2, d orbitals on the metal ion.
-
+
The Journal of Physical Chemistry
The importance of T bonding to kinetics is in its effect on A, the crystal field stabilization energy. The substitution rate constant kon is correlated with the loss of crystal field stabilization energy, being smaller the larger the energy loss in attaining the activated complex structure. Therefore, the enhancement may arise from this effect when A is decreased by n-orbital interaction. Such is likely the situation for picolinate and terpyridine.’O It might be mentioned that it is quite possible that in other cases the interaction may be of small consequence, bearing relatively little effect on the enhancement or that the n interaction may even increase the A value, thus diminishing the rate of water loss. The relative role played by T bonding and back T bonding in the metal chelate will determine the effect on AI1 and hence on the rate constants.
Acknowledgment. The authors gratefully acknowledge partial support from Public Health Service Research Grant GM-08893-06 from the National Institute of General Medical Sciences, Public Health Service, from the National Science Foundation for College Teacher Research Participation Grant XO. GY3796 (A. K.), and from the Petroleum Research Fund for Grant 2982B (R. F’. P.). (8) (a) K. Kustin, R. F. Pasternack, and E. M . Weinstoclr, J . Amer. Chem. Sac., 88, 4610 (1966); (b) A. Kowalak, K. Kustin, R. F. Pasternack, and 8. Petrucci, ibid., 89, 3126 (1967). (9) F. Basolo and R. G. Pearson, “Mechanisms in Inorganic Reactions,” 2nd ed, John Wiley & Sons, Inc., New Yorlr, N. Y., 1967, p 147.
(10) R. H. Holyer, C. D. Hubbard, 8. F. A. Kettle, and R. G. Wilkins, Inorg. Chem., 5 , 622 (1966). (11) 3’. A. Cotton and G. Wilkinson, “Advanced Inorganic Chemiatry,” 2nd ed, Interscience Publishers, Inc., New York, N. Y., 1966, p 708.
The Photoxidation of t-Butyl Iodide at Low Temperatures
by Theodore Mill and Roger Stringham Department of Physical Organic Chemistry, Stanford Research Institute, Menlo P a r k , California 94086 (Received August 68, 1968)
We have examined the photolysis of t-butyl iodide in oxygen-saturated CFCL around - 100” as a possibly efficient source of t-butyl peroxy radicals for synthesis of di-t-butyl tetroxide and trioxide.‘ At 25” t-butyl iodide is reported to give t-butyl radicals through photolysis of the C-I bondU2 Disproportionation of the t-butyl radicals gives isobutane and isobutylene, the (1) T. Mill and R. Stringham, J, A m e r . Chem. Sac., 90,1062 (1968). (2) C. E. McCawley, W. H. Hamill, and R. R. Williams, ibid., 76, 6263(1864).
283
NOTES major products of the reaction. The following reaction scheme is expected on the basis of our earlier work with t-butyl radicals and oxygen' hv
t-Bu.
+ 1.
(1)
0 2 --t
t-BuO,
(2)
t-BuI t-Bu.
+
__
2t-Bu02 t-BU204 > -70' t-BuzOi 2t-BuO. 02 t-BuO
- + L-BuOZ
(3)
+
