NOTES
2018
normal for a unimolecular decomposition and is quite similar to those of other retro-Diels-Alder reactions. * I n fact, this thermal decomposition is so “normal” hat very little mechanistic delineation is possible. (7) N. E. Duncan and G. J. J a m , J . Chem. Phys., 20, 1644 (1952). (8) C. Walling, “Free Radicals in Solution,” John Wiley and Sons, Inc., New York, N. Y., 1957, Chapter 2.
Free-Radical Substitution in Aliphatic Compounds. VI.
The Reaction of Fluorine Atoms with Carbon Tetrachloride1
by D. T. Clark and J. JI. Tedder Department of Chemistry, T h e University, Shefield 10, England (Received January $1, 1964)
I
I
1.4
1.5 1.6 I/T x io3
1.7
Figure 1. Arrhenius plot for bicyclo[2.2.1]hept-2-ene pyrolysis, 0 (stirred flow reactor), 0 (tubular flow reactor).
The activation energy found in this reaction (42.8 kcal.) is intermediate between that found for dicyclopentadiene (34.2 kcal.)*” and that found for bicyclo[2.2.l]hepta-2,5-diene (50.2 kcal.).’ It differs considerably from those reported for cyclohexene pyrolysis (57.s4d or 67.64h kcal.) and from that calculated for 4-vinylcyclohexene pyrolysis (-62 kcal.) Semiquantitatively, the values for the bicyclic compounds support a rate-determining step in which only one bond i s broken. Bond breaking in dicyclopentadiene procluces two allyl radicals, in bicyclohepteiie an allyl and alkyl radical, and in bicycloheptadiene an allyl and irinyl radical. The activation eiiergies for these processes8 should, therefore, lie in the order which is t‘ound; however, the established order does not eliminate the possibility of a mechanism in which two bonds are broken simultaneously. The pre-exponential factor for the bicyclo [2.2.1]hept-Bene pyrolysis ‘is of the magnitude considered
.’
The Journal of Physical Chemistry
Previous work in this laboratory has been devoted to the study of the abstraction of hydrogen from substituted aliphatic hydrocarbons by halogen atoms2 or trichloromethyl radicals. A fairly complete picture of this process has been d e v e l ~ p e d . ~The present investigation is the first of a series in which the abstraction of halogen atoms from aliphatic compounds will be studied. The fluorination of carbon tetrachloride has been studied qualitatively by Ruff5 and by Simons.6 Ruff found little reaction a t room temperature; and when the carbon tetrachloride was refluxed, explosions occurred. Both Ruff and (subsequently) Simons carried out the heterogeneous fluorination of carbon tetrachloride, bubbling the gaseous fluorine through refluxing carbon tetrachloride and adding various “catalysts” such as iodine, arsenic, or bromine. Simons also studied the fluorination of difluorodichloromethane in the gas phase, again in the presence of “catalysts.” The notable feature of this work was that difluorodichloromethane was much less reactive than carbon ~
~~
~~
~
(1) This research was supported in part by Aeronautical Systems Division, AfSC. through the European Office, Aerospace Research, U.S.A.F., Grant No. AF.EOARDC-61-7. (2) P. C. Anson, P. S. Fredricks, and J. M. Tedder, J . Chem. SOC., 918 (1959): P. S. Fredrioks and J. M. Tedder, ibid., 144 (1960); ibid., 3520 (1961); I. Galiba, J. M. Tedder, and R. A. Watson, ibid., in press. (3) B. P. McGrath and J. M .Tedder, Bull. SOC.Chim. Belges, 71, 772 (1962). (4) J. M. Tedder, Quart. Rea. (London), 14, 336 (1960). (5) 0. Ruff and R. Keim, 2. anorg. allgem. Chem., 201, 245 (1931). (6) J. H. Simons, R. L. Bond, and R. E. McArthur, J . Am. Chem. SOC.,62, 3477 (1940).
NOTES
2019
tetrachloride. The present work was planned in the light of these qualitative reports. However in the event they proved rather misleading. If fluorine and carbon tetrachloride vapor react together the followirig steps represent the expected reaction paths
Fz +2F’
+ CCl, +Cc13‘ + FC1 cc13’ + Fz -+ CC1,F + F’
F’
(1) (2)
(3)
The possible chain-terminating steps are as follows
+ F‘ +Fz CC13’ + F’ + CCldF F“
cc13’ f C c & ’--+ CzC&
(4) (5)
(6)
The reaction chain contains no branches, so provided there is no thermal branching, the concentration of radicals cannot exceed the concentration of fluorine atoms in equilibrium at the reaction temperature. As this is very small, we can neglect reaction 6. Reaction 3 is exothermic to the order of 60 kcal./mole and will therefore have no appreciable activation energy. At room temperature the concentration of fluorine atoms is approximately times the concentration of molecular fluorine, and the concentration of trichloromethyl radicals must certainly be much less than this. Therefore, provided the concentration of fluorine is at least as great as the concentration of carbon tetrachloride the rate of reaction 5 must be vanishingly small since all the trichloromethyl radicals will react by reactioin 3 before they have a chance of meeting a fluorine atom. (In practice a tenfold excess of fluorine has been employed in the present work.) This leaves reaction 4, L e . , the recombination of fluorine atoms, as the only chain-terminating step. IF this argument is correct we would expect very long chains and this is in accord with the experimental results. We, thus, have an exceptionally simple system and making the usual steady assumptions we obtain the following expression for the rate of reaction 2. d[FCl] -~ =
dt
IC,
(2)
‘/a
[F2]”’[CCL]
The ratio lcl/ke iEi the dissociation constant of fluorine, K , and since the reaction chain neither increases or decreases .the number of fluorine atoms, we are quite justified in using the value obtained from equilibrium studies.
Experimental Fluorine was used directly from a generator, but the’ 10-amp. cell was run at about 1 amp. for 24 hr. and then at maximum current for a further 3 hr. before each reaction run. The exclusion of oxygen from the fluorine was vital and in view of the consistency of the results we believe that oxygen had been eliminated. I n the initial experiments when the cell was only run a t 3 amp. for 1 hr., the reaction was very much slower and the results were erratic. The fluorine passed through a trap at -70’ to remove hydrogen fluoride and was then joined by a stream of (‘oxygenfree” nitrogen. The carbon tetrachloride was vaporized by bubbling a stream of nitrogen through the liquid which was maintained a t a constant temperature. The reaction vessel consisted simply of a T-piece in which the carbon tetrachloride and fluorine streams met at right angles. The T-piece and the tubes leading into it were immersed in a thermostat. The combined flow passed along 3 cm. of 3/16 in. 0.d. copper tubing still in the thermostat before being led directly into a trap a t --70’. Unchanged carbon tetrachloride and trichlorofluoromethane condensed in the trap while the fluorine monochloride and unreacted fluorine passed on to a heated column into which excess hydrogen was also led. The reaction occurred as an actual flame and the hydrogen fluoride and hydrogen chloride were collected in a series of traps cooled by liquid nitrogen. When the run was completed, heated hydrogen was passed through the traps and then through 2 N sodium hydroxide solution free from carbonate (it was not possible to pass the reaction gases directly through alkali because the height of liquid required to ensure complete absorption of the hydrogen halides caused too much back pressure). The fluoride was determined by acidification of a measured portion of the solution followed by the addition of excess of a standard calcium solution and back titration with standard EDTA. The chloride was determined spectrophotometrically by a modification of the method of Kitano and Tsubota.7 Six runs at 20’ are tabulated in Table I. Each run was of 20 min. (i10 sec.) duration and the volume ( V ) of the reactor was 5 cc. Since the reaction only went to 10% or less a plot of carbon tetrachloride concentration against ( [ F C ~ ] U ” / ” [ F Z ] -X~ ” lo2) gave a straight line with a slope equal to l/K”/”kzVX 60. See Fig. 1.
Discussion The fact that the experimental results lie on a straight line gives strong support to the rate equation Volume 68, Number 7
J u l y , 1964
2020
NOTES
Dissociation Energies of Alkaline E a r t h
Table I
Oxides' Concentration in moles/min. Flou.rate Fz X 10' CClr X 106 FC1 X 108 ( U ) , co./min.
13 27
0 98
16 96 18 46 15 I 1
1 2 2 25
3 28 3 50 4 38
11 80 14 67
14 5 20 2 31 9 43 5
45 1
[FClI X W e X 10* [Fzl'/2
213 213 225
3 91
22 I
11 77 14 21 17 87
65 0
Brookhaven National Laboratoru, Upton, N e w York (Received February 5, 1964)
4 83 7 93
227 223
discussed above. The slope of the line gives k2 = 13.7 K - ' / E cc. mole-' sec.-' and calculating K from the data of Cole, I;arber, and Elverurn* gives k2293= 2X molesu1 cc. see.-'. This is a very reasonable result and if the pre-exponential factor is of the order of l O I 5 this gives an activation energy of 2.5-3 kcal./
/
Gradient
, CClJl16 mdp 0
I
2
3
- 3x10'
12.3 x IOa2
r d
by M. S. Chandrasekharaiah
Presently available dissociation energies for gaseous alkaline earth oxide molecules may not be the true dissociation energies of the ground molecular states. Dissociation energies obtained from spectroscopic analysis appear to be for some excited state.2a A reliable thermochemical evaluation is not possible due to the lack of accurate vapor pressure measurements. Only for BaO are there any reliable vaporization datazbthat can be applied in a good estimate of D ofor BaO(g) through a suitable thermochemical cycle. An attempt to estimate the dissociation energies based on a simple ionic niodel is made in this work. An ionic model consisting of ?\I+zion and 0-2 ion separated by the equilibrium internuclear distance is assumed for the molecules, and the binding energies are calculated. The binding energy of such a system is the electrostatic interaction energy between the cation and the anion at the equilibrium internuclear distance of separation. The expression for such interaction energy can he derived from the well-established relations of electrostatics.3 Seglecting dispersion energy and zero point energy terms, the following expressions can be derived for the potential energy of such a system.
I
4
5
Figure 1. Plot of carbon tetrachloride concentration against [FCl]Ua/%[F2] -'h X lo2.
mole which is also of the expected order. Experiments were also carried out at 309.5OK. and the rate of reaction increased to such an extent that the present method of collection and analysis proved inadequate. Insofar that any conclusions could be drawn from the results at 309.5OK. they confirmed the present mechanism giving an over-a11 activation energy (Ez 0.5AN) of greater than 20 kcal. mole-'.
+
(7) H. Kitanoand H. Tsubota, J . Chela. SOCJ a p a n , 75a, 931 (1954). (8) L. G. Cole, M. Farber, and G. W. Elverurn. Jr., J . Chem. Phys., 20,
586 (1953).
T h e Journal of Physical Chemistry
where W is the potential energy of the system, a1 and cy2 are the polarizabilities, G~ and p2 are the induced dipole moments of the cations and anion, respectively, e is the electronic charge, A and p are the repulsive potential parameters, and r is the distance of separa(1) This work was performed under the auspices of the U. S. Atomic Energy Commission. (2) (a) L. Brewer, Chem. Ret., 53, 1 (1953): (b) R. J. Ackermann and R. J. Thorn, Progi. Ceramic Sci., 1, 39 (1961). (3) E. A. Moelwyn-Hughes, "Physical Chemistry," Pergamon Press, New York, N. Y., 1957, pp. 292-333.
