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would have to assume that the OCH, group attached to the carbonyl carbon would reverse the polarity of the carbonyl bond by induction if other factors were not important, which seems to be rather unlikely. In this type of calculation, however, we have so far tacitly assumed that the average energy of excitation, AE, which appears in the denominator of Pople’s expression for paramagnetic shielding5 is practically constant in a homologous series of molecules. This may not and need not be actually so, and as little as 10% uncertainty in AE may lead to about 40-p.p.m. uncertainty in the carbonyl shift,s. Furthermore, the constitutive corrections in these systems may reflect not only the interaction of substituents with the ?r-systems but also an additional effect of changing from sp3 to sp2 hybridization. The effect of resonance on the carbon shifts in these systems is very difficult to assess even on a qualitative basis. It must be borne in mind that in the contributing resonarice structures, such as 0-
the carbons in question are merely resonance transmitting. They cannot be qualitatively compared with ortho and para carbons in substituted benzenes on which the development of charge can be discussed in terms of enhanced screening. On the other hand, resonance may contribute to a more asymmetric charge distribution around the carbons in question and hence to increased anisotropy of their electronic environment. Another feature of resonance which cannot be ignored is the extent to which it may affect the average excitation energy, AE, in these molecules. Thus, according to Jaffd and Orchin,e substitution of a group with a lone pair as NR2, OR, or halogen, into the carbonyl group, as in amide, ester, or acid chloride results in a pronounced effect on the n + P* transition energy. The atoms with the lone-pair electrons influence such transitions both by resonance and inductive effect. Even in the absence of inductive effect the P* level is raised by resonance, thus leading to higher excitation energies. The lone-pair atoms, 0, N, and C1, are more electronegative than carbon and hydrogen, and the lone-pair electrons on the carbonyl oxygen are held more firmly than they would be in ketones and aldehydes. The effects of resonance and induction are difficult to separate, but the n + a* excitation energies follow the order 0 > N > C1 > H in CH3COX compounds, where 0, N, C1, and H are the atoms of the group X immediately attached to the carbonyl. It is noteworthy that the deviations in carbonyl shifts in The Journal c ~ f Physical Chemistry
ester, amide, and acid chloride with respect to acetaldehyde follow exactly the same trend. If n + ?r* transition energies are to be assumed to follow the same order as the average excitation energies, AE, then this would be in agreement with Pople’s f ~ r m u l a t i o n , ~ according to which high excitation energies lead to high-field shifts. The very range of variation in n + P* transition energy (it is about 30Oj, higher in esters than in acetaldehyde) accounts for the right order of magnitude of the deviations from constitutive additivity in these molecules. It would thus be tempting to conclude that, at least in the case of carbonyl shifts, the energy of excitation, AE, plays the dominant role. However, in the case of the benzene series such an explanation is untenable, since substitution by X usually leads to the lowering in AE. If other factors in the CsH6X series are more important, the empirical correlation (2) will have to be regarded, a t best, as being merely fortuitous. We have thus been unable to find a simple and unambiguous explanation for the trends in deviations from constitutive additivity in these two systems. We feel, however, that the statement of the problem and the factors involved, conflicting as they seem to be a t present, may help in the future attempts to understand the delicate interplay of various effects which contribute to the over-all C13 chemical shifts in substituted hydrocarbons.
Acknowledgment. The authors wish to express their gratitude to the Research Corp. for the financial support of this work and to Professor G. E. Maciel for helpful discussions and suggestions. (6) H. H. Jeff6 and M.Orchin, “Theory and Applications of Ultraviolet Spectroscopy,” John Wiley and Sons, Inc., New York, N. Y . , 1962, pp. 178-181.
Carbon-13 Isotope Effects in the Oxidation of Acetic Acid
by Mieczyslaw Zieliliski’ Enrico Fermi Institute f o r Nuclear Studies, University of Chicago, Chicago, Illinois (Received October 6, 1964)
The oxidation of organic compounds by the Van Slyke-Folch2 “wet” method is a convenient and widely applied procedure. Acetic acid, because of its crit,ical importance in metabolism, is often a sample material
NOTES
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for such oxidation and, also frequently, isotopically labeled acetic acid or acetate is the isolate under study. Earlier investigations3v4 have demonstrated that there is an appreciable kinetic isotope effect associated with the wet oxidation of acetic acid (as acetate) labeled with C14; the results, however, are discrepant. Evans and Huston3 do not indicate the degree or increment of decomposition involved in their experiments (in which specific activity measurements were made on “two consecutive fractions of roughly the same weight”), but Eyring and Cagle5 estimated their results to be equivalent to a 5% isotope effect in the oxidation of the carboxyl-labeled acetate. In the later experiment^,^ an effect of 1.7 f 0.5% was observed in the oxidation of carboxyl-labeled acetate, but a 4.0 f 0.6% effect in that of the methyl-labeled compound was observed. Because of the intrinsically more accurate measurements possible of CI3 isotope effects compared with those of CI4,it was of interest to attempt to resolve the discrepancy described above by investigations of the oxidation of acetate containing CI3 a t the natural abundance level.
From the three experiments we obtain k l / k z = 1.0098 f 0.0030. Combination with the assumed value for (k3/kZ) yields k l / k 3 = 1.0221 f 0.005. Though the value of kl/kr is not independent of the earlier CI4 results, it is, nevertheless, of interest to compare these isotopic rate constant ratios for the C’3 effects with those calculated directly from the CI4 data. We have ( k l / k 2 ) e a l c d = 1.0087 (compare with 1.0098) and ( k l / k 3 ) c a l c d = 1.021 (compare with 1.0221). The present results are thus consistent with the earlier findings of Zlotowski and Zielinski for the wet oxidation of acetic acid.
Experimental
(1) Department of Nuclear Chemistry, University of Warsaw, Poland. (2) D. D. Van Slyke and J. Folch, J . Biol. Chem., 136, 509 (1940). (3) E.A. Evans and J. L. Huston, J . Chem. Phya., 19, 1214 (1951). (4) I. Zlotowski and M. Zielinski, Nukleonika, 4, 5 (1959). (5) H. Eyring and F. W. Cagle, J. Phys. Chem., 56, 889 (1952).
AcknowZedgment. The author expresses his thanks to Professor Robert N. Clayton for making possible his stay at the Enrico Fermi Institute and for assistance in the preparation of this note. Also, the author expresses appreciation to Professor Peter E. Yankwich, the referee of this note, for introducing substantial changes which made the whole note clearer and more systematic. The research was supported in part by a grant from the National Science Foundation (NSF-GP-2019).
Samples of acetate mixed with excess solid sodium hydroxide were pyrolyzed quantitatively in copper vessels to methane and carbonate-the former being oxidized to carbon dioxide over copper oxide, the latter being converted to carbon dioxide with phosphoric acid. Isotope ratio measurements on these samples permitted evaluation of the ratio (C13HaC1zOOH/C12H2C1300H) Magnetic Resonance of the Triplet State of = Q. Qowas found to be 1.0074for the acetate used, and the average of the methane-carbon and carboxyl-carOriented Pyrene Moleculesla bon isotope ratios was identical with that obtained for carbon dioxide resulting from complete combustion by 0. Hayes Griffithlb (over copper oxide) of the original acetate, Ro. The partial oxidation experiments were carried out Gates and Crellin Laboratories of Chemistry, lo California Institute of Technology, Paaadena, California (Received November 2, 1964) using the Van Slyke-Folch procedure. Values of Qf and R , (the ratio indicated above and the ratio Cl3O2/ CI2O2at time 1 or degree of reaction f) were obtained Bree and Vilkos2 have recently completed an optical in each of three runs. polarization study of the lower singlet states of pyrene Where kl is the specific rate constant for the oxidain a fluorene matrix a t 77°K. The crystal structure of tion of C12H3C1200H,k2 that for C12HaC1300H,and fluorene is orthorhombic, the space group is Pnam, k3 that for Cl3H3ClZ00H,it can be shown that and there are four molecules per unit cell.3 This crystal is a convenient matrix because the long axes of the fluorene molecules are all parallel to the crystalline c axis. We report here the observation of electron Use of this formula with the present data requires a value for k3/k2; this we take from the CI4 results of Zlotowski and Zielinski to be 0.9879 f 0.004, assuming that the C14 kinetic isotope effect would be twice the C13 effect.
(1) (a) Supported by the National Science Foundation under Grant No. GP-930; (b) National Sbience Foundation Predoctoral Fellow; (c) Contribution No. 3181. (2) A. Bree and V . V. B. Vilkos, J . Chem. Phys., 40, 3125 (1964). (3) D.M.Burns and J. Iball, Proc. Roy. SOC.(London), A227, 200 (1955).
Volume 69,Number 4
April 1065
