NOTES
664
tion is explained immediately by eq. 5: I n pure ethanol the limited dissociation of nitric acid, together with the side reaction producing ethyl nitrate rather than nitric acid to the extent of -30%, makes the virtual constancy of k“ more difficult to explain, though the ionization of the nitric acid is still sufficient to make the reaction difficult to follow conductometrically.13 I n the case of nonhydroxylic solvents, such as acetonitrile, where no nitric acid is produced, the value of IC!’ does, as expected, decrease during the course of reaction.8 Furthermore, when silver lactate replaces silver nitrat’e in ethanol-water mixtures, the value of k“ again decreases as the reaction proceeds14which is exactly as expected in view of the weak nature of lactic acid. The ion-pair mechanism predicts that the reaction rate will decrease as the dielectric constant is raised in moving from ethanol to water as solvent, by the ratio of the (1 - y ) values in the two solvents which is 4 3 0 ( y in water calculated from an extrapolated KA valuelo and stoichiometric activity coefficient^'^). As observed, the reaction with methyl or ethyl iodide3 is only decreased by a factor of -7; the rate does not change monotonously, but it is riot expected that the silver nitrate ion pair will remain equally reactive in conditions of competitive solvation obtaining in niixtures of polar solvents. The recent work of ;\Ielendez-Andreu5 using an acetone-water mixture provides good support for the proposed mechanism involving the silver nitrate ion pair. The reactions of alkyl bromides, from methyl to nhexyl, with silver nitrate and perchlorate were studied and found (by means of the Leffler ploti5) to be of the same mechanism for all the bromides in all the solvent mixtures. The reaction with silver nitrate is particularly rapid in acetone, and it is significant that silver nitrate is a particularly weak electrolyte in this solvent. Silver perchlorate, which is not as weak in acetone,16 reacts niuch more slowly, even though in water-rich mixtures, where the two electrolytes are both strong, the rates are very similar. It is evident that the rapid rate of the silver nitrate reaction in acetone cannot be due to a high reactivity of the silver ion in this solvent and also can hardly be due to the increased nucleophilic character of a desolvated nitrate ion since the transport number for this ion in silver nitrate in acetone (0.58) is little different from that in water (0.54) and less than that in ethanol (0.61). The considerable association of silver nitrate in acetone, which may be due to stabilization of the ion pair itself in this solvent , I 7 explains the high reaction rate, at once, if the 1011 pair is the reacting species. We conclude, therefore, that thc suggestion of The Journal of Physical Chemistry
Donnan as to the mechanism of the silver nitrate-alkyl halide reaction was essent,ially correct for many solvents despite later criticisms. (13) A. G. Walton, Thesis, University of Nottingham, 1960. (14) R. Parsons, “Handbook of Electrochemical Constants,” Butterworth and Co., Ltd., London, 1959. (15) J. E.Leffler, J . Org. C h a . , 20, 1202 (1955). (16) V. S. Griffiths, K. S. Lawrence, and AM.L. Pearce, J . Chem. Soc., 3998 (1958). (17) W. R.Gilkerson, J . Chem. Phy8., 2 5 , 1199 (1956).
The Reaction of Methyl Radicals with Methyl and Methylene Fluoride by G. 0. Pritchard, J. T . Bryant, and R. L. Thommarson Department of Chemistry, University of California, Santa Barbara, California 93018 (Received July 29, 1964)
Raal and Steacie‘ have observed a significant decrease in the activation energy for H atom abstraction by methyl radicals with increasing halogenation of methane
CHs
+ CH, - .X,
+CH4
+ CHa - ,Xn
(1)
For X = chlorine, and n = 1, 2, and 3, they obtained values of El of 9.4, 7.2, and 5.8 kcal. mole-’, respectively. For X = fluorine, and n = 1 and 2, El was found to decrease from 8.7 to 6.2 kcal. mole-’. A similar trend was obtained with methyl and methylene bromide. They conclude that there is a progressive decline in the magnitude of the C-H bond strength with increasing halogenation, leading to an enhanced “activity” of hydrogen atoms in substituted methanes. While this generalization is correct for chloromethanes,2 it is certainly not correct for fluoromethanes.2 We3 have found El using CF,H to be 10.2 f- 0.2 kcal. mole-’, and the most likely value314for D(CF3-H) is close to 105 kcal. mole-’. We have therefore redetermined El for CFHl and CF2H2. (1) F.A. Raal and E. W. R. Steacie, J . Chem. Phys., 20, 578 (1952). (2) C . T. Mortimer, “Reaction Heats and Bond Strengths,” I’ergamon Press, London, 1962,pp. 132-134. (3) G. 0.Pritchard and R. L. Thommarson, J . Phys. Chem., 68, 568 (1964). (4) E.Whittle, private communication: from unpublished bromina tion experiments. This is some 4 kcal. lower than the previous value, see 1’. Corbett, A. h f . Tarr. and E. Whittle, Trans. Faraday Soc., 59, 1609 (1963).
NOTES
665
The experimental method was identical with that described previously for the CFIH experiment^,^ in which reaction 1 competed with CD3
+ CD3COCD3 -+- CD, + CDzCOCDs
(2)
The fluoromethanes were obtained commercially and purified by repeated low temperature fractionation. Their mass spectra indicated no detectable impurities and they were coincident with the previously published spectra for the two compounds.j In the CFH3 experiments, the pressure of CFH3 was approximately 8 cm., and that of the ketone about 3 cm. I n the CF,H2 experiments, the respective pressures were 5 and 3.5 em. The data for CFzHzare collected in Table I. The CD3H/CD4ratios have been corrected for the ds impurity in the acetone-ds.
Table 1: Data for the Competitive System: CHtF2 +CD3H CHFI (1) and CDI CD3 CDaCOCDa --+ CDa CDzCOCD3 (2)
+
+
+ +
[acetone-dsl
[CDaHl [CD41
ki -
O C .
122 143 164 1'77 198 222 239 268 340'
1,03 1.72 1,69 1,48 1.12 1.73 1.55 2.11 1.36
1.22 1.64 1.52 1.29 0.98 1.30 1.08 1.29 0.82
1.19 1,29 1.37 1.44 1.47 1.57 1.70 1.83 1.91
Temp.,
[CHzFzI -
ki
A least-squares treatment of the Arrhenius plot for CF2H:! gives k,/kl = 5.5 exp[-l200/RT]. Taking E , = 11.4 f 0.2 kcal. mole-' (assuming zero activation energy for CD, radical recombination) , we obtain El (for CFzH,) = 10.2 f 0.2 ltcal. mole-'. An average of 17 runs with CFH3 in the temperature range 193 to 331' yielded k z / k l = 3.2 f 0.2 (the ratio being 3.18 a t 193' and 3.41 a t 331'). This scatter was not particularly satisfactory, but it seems valid to equate El (for CFH,) with E, = 11.4 f 0.2 kcal. mole-'. I t is apparent that the activation energies given by Raal arid Steacie' for these two reactions are markedly in error. Trotman-Dickenson" has remarked that the results may be unreliable because coinplicating side reactions probably occurred. The differences in the two sets of determinations are well beyond the limit of being low by 1 kcal. mole-', an effect which has been observed when acetone-do is used as the radical source.'
However, we must comment that below 200' we were unable to obtain reproducible data with the CFH3 system. This was despite repeated careful checks on the purity of the compounds and the condition of the quartz reaction cell. Periodic blank runs with acetone-& alone consistently gave reproducible W 3 H / CD, ratios. From 200' down to -100' we obtairted generally decreasing values of k z / k l , which were very scattered. These indicated that E , - E1 = -2.7 f 0.5 kcal. mole-', which yields E , = 8.7 f 0.7 ltcal. mole-', in complete, but probably fortuitous, agreement with Raal and Steacie's figure' for this reaction. These authors also obtained their result in the temperature range 125 to 211'. Our data from the CF2Hz system were completely reproducible over the entire temperature range, and a comparison of the results derived from the three fluoromethane systems strongly suggests that we disregard the data obtained below 200' in the CFHa system. The cause of the discrepancy is not obvious; it is plausible that an undetected trace impurity of a substance containing very labile H atoms is responsible. Taking the value for the pre-exponential factor6 A, = 6.3 X 10'' mole-' cc. see.-' gives normal steric factors for the two reactions of about We intend to present data on reaction -1 for CFH, and CFzH radicals. Our initial work* on the photolysis of (CFH2),C0 indicates that the reactivity of CFHz radicals with regard to H atom abstraction lies between that of CF3 and CH3 radicals. The activation energies for the two reactions
+ CH, +CF3H + CH3 CH3 + CH, --+- CH, + CH3
CF3
are about 10 and 14 Itcal. niole-l, re~pectively.~We may therefore postulate that reaction 1 for CI"H3 and CFzHz will be close to being thermoneutral, and that D(CFH2-H) and D(CF,H-H) will not be very different from D(CH3-H). Aclrnowledgnzents. We are very grateful to Dr. Whittle for communication of his results prior to publication, and to the Sational Science Foundation for support of this work. (5) ,J. I{. XIajer, Adcan. Flmrine Chem., 2 , 55 (1961). (8) A. F. Trotman-Dickenson, "Gas Kinetics," Butterworth arid Co.. Ltd., London, 1955, p . 202. (7) H. 0 . I'ritchard and G . 0. I'ritchard, Can. .I. Chrm., 41, 3042 (1963). ( 8 ) G. 0. I'ritchard, 11. Venugopalan, and T. 1.; Graham. .I. P h y s . Chem., 68, 1786 (1964). (9) G . 0. I'ritchard and C.. H. Miller, .I. C'hcm. P h l j s . . 35, 1135 (1961).
Volrime 6.9. Number 2
F?hr?inry 1966
