NOTES
1952
measured. For group D alloys there may be anticipated to be large differences in the solubilities at certain added metal contents. These predictions have been verified for only two alloy systems: the platinumpalladium2 and the gold-palladium (ref 4 and this work).
Acknowledgments. This research was supported by the U. S. Atomic Energy Commission. The authors are most appreciative of this financial support. The authors are also indebted to Engelhard Industries, Inc., for the gold-palladium alloys used in this research.
Reactions of Diatomic Molecules. IV. Kinetics of Formation of Bromine Chloride by Peter R. Walton and Richard M. Noyea Department of Chemistry, university of Oregon, Eugene, Oregon 97409 (Received December 6 , 1966)
Calculations described elsewhere’ predict that all reactions of halogens with halogens will proceed by bimolecular mechanisms. For the formation of BrCl from the elements, the predicted activation energy is about 13 kcal/mole. Previous studies of this reaction are mostly qualitative. Jost2 and Brauer and VictorS found rapid reaction in gas phase, and Jost2 even estimated an activation energy of 14 kcal/mole. These studies did not demonstrate homogeneity of the reaction, and surface effects are obviously hard to eliminate. Although reactions in solution are more easily shown to be homogeneous, “inertness” of solvent can never be demonstrated unequivocally. Barratt and Stein4 observed the reaction to be “instantaneous” in ether and chloroform but to have a time lag of several seconds in carbon tetrachloride. Dennis Forbess at the University of Oregon also observed a measurable rate in this solvent, but he could not get reproducible results. Our own qualitative observations indicated that the rate in carbon tetrachloride was strongly accelerated by traces of moisture as reported by Hildebrands for the formation of IC1 in this solvent. This sensitivity to moisture suggested the use of pure sulfuric acid as a reaction medium. Visual examination indicated that the spectra of the halogens were the same in this solvent as in gas phase. Absorption spectra of bromine, chlorine, and a reacted mixture agreed well with those reported at wavelengths longer than 3100 A in carbon tetrachloride by Popov and Mannion,6 The Journal of .Physical Chemistry
and the position of equilibrium did not appear to be shifted from their observations. The rate of reaction was much slower than in organic solvents, was satisfactorily reproducible, and was not affected by the deliberate addition of small amounts of water. These facts offered enough evidence of “inertness” to justify the kinetic observations reported here.
Experimental Section Materials. Reagent grade bromine was purified by shaking with concentrated sulfuric acid followed by distillation. Chlorine was purified by passing through concentrated sulfuric acid. The solvent was 99.6% sulfuric acid prepared by adding reagent grade oleum to 98% acid. The composition was analyzed by electrical conductivity. Solutions of the halogens were made up prior to each run. Since air-saturated solutions of chlorine were found to undergo photochemical deterioration when heated, the solvent used in preparing solutions had been saturated with dry nitrogen. The solutions were analyzed by diluting with ice water containing potassium iodide and then titrating with standardized thiosulfate. Procedure. Mixtures of the desired composition were prepared directly in optical cells and thermostated to the desired temperature. Optical absorbance was then followed as a function of time with a Beckman spectrophotometer. The observations could be fitted satisfactorily to the equation for a reversible bimolecular process, and a computer program with a variable infinite time absorbance was used to get the best fit and to compute the rate constant. The calculations also required information as to the position of equilibrium. The equilibrium constant for the reaction
Brz
+ Cla
2BrC1
(1)
has been estimated to be almost independent of temperature and to be about 7 in vapor phase7 and in carbon tetrachloride? Very rough measurements at 75.6’ confirmed that the value in sulfuric acid is approximately the same or slightly less. Since uncertainties in this quantity have very little effect on (1) R. M. Noyes, J . Am. Chem. SOC.,88, 4318 (1966). (2) W. Jost, 2.Phyeik. Chem., B14, 413 (1931). (3) G. Brauer and E. Victor, Z . Elektrochm., 41, 508 (1935). (4) S. Barratt and C. P. Stein, Proc. Roy. SOC.(London), A122, 682 (1929). (5) J. H. Hildebrand, J . Am. Chem. SOC.,68, 916 (1946). (6) A. I. Popov and J. J. Mannion, ibid., 74, 222 (1952). (7) C. M. Beeson and D. M. Yost, ibid., 61, 1432 (1939).
NOTES
1953
rate constants measured when one reactant is in considerable excess, no attempt was made to measure the equilibrium constant more accurately.
Results Since bromine absorbs much more strongly than chlorine at wavelengths of analytical interest, the absorbance change during a run was greatest when chlorine was the species in excess, and most runs were made for this condition. An extensive series of such runs at 75.6’ supported the hypothesis that the rate was indeed first order in each reactant, and the same rate constant was computed for one run in which bromine was in twofold excess over chlorine. For a series of runs at the same temperature in which chlorine and bromine were initially in equal (and stoichiometric) concentrations, the computed rate constants varied approximately inversely with initial concentration. Such behavior implies that the over-all order is little more than 1 instead of the anticipated value of 2. At several different temperatures, the rate constant for a solution 0.06 M in chlorine and 0.005 M in bromine was over twice that for a solution 0.04 M in each halogen. We are unable to explain these kinetic anomalies or to derive a rate expression consistent with all of the observations. The runs in which one reactant was in excess definitely support the kinetics anticipated for a bimolecular mechanism, and the alternative atomic chain mechanism is at least as unsatisfactory with regard to kinetic consistency and is also untenable for reasons discussed below. The reported rate constants are therefore restricted to runs at unequal reactant concentrations. Let k be the rate constant for reaction 1 from left to right. Estimated values of k are presented in Table I. The data can be fitted by the equation log k =
Discussion If the solvent is truly “inert,” the reaction presumably proceeds either by a bimolecular mechanism or by an atomic chain initiated by dissociation of bromine.* If the atomic chain mechanism applies, the apparent activation energy can hardly be significantly less than ‘/&‘(Brz) Doo(Clz)- Do’(BrC1)= 27.7 kcal/mole. The very much smaller value observed makes the chain mechanism impossible. Further support for this conclusion was provided by a run at 75.6’ that was illuminated with a conventional incandescent lamp. The rate was not significantly increased, although the concentration of free bromine atoms must have been at least one order of magnitude greater than in the thermal runs. Because of the compelling evidence against the atomic chain mechanism, we believe that even the crude kinetic measurements reported here are sufficient to demonstrate the bimolecular mechanism. The only alternative would seem to be the intervention of some ionic species in this highly polar medium, and we are not aware of any evidence for plausible concentrations of such species. The observed activation energy of 17 kcal/mole is in gratifying agreement with the 13 kcal/mole predicted’ for the gas phase; it is certainly very much less than the 52.9 kcal/mole predicted by Benson and Haugens for the same reaction.
+
Acknowledgment. This work was supported in part by a grant from the U. S. Army Research Office (Durham). (8) R. M. Noyes, J. Am. Chem. SOC.,88, 4311 (1966). (9) 5. W.Benson and G. R.Haugen, ibid., 87,4036 (1965).
The Yield of Thermal Ethyl Radicals from the Table I : Rate Constants for Reaction 1
Radiolysis of Ethylene-Cyclohexane Solutions
Temp,
104k, l./mole
OC
sec
50.1 59.8 67.4 75.6 84.7
7.5 15 24 45 100
8.3 - 17.0/8, where 8 is 2.303RT in kilocalories per mole. Although the activation energy is probably uncertain by about 2 kcal/mole, these results are sufficiently precise to give the desired mechanistic information.
by J. L. McCrumb and Robert H. Schuler Radiotion Research Laboratories, Mellon Institute, Pittsbu~gh,Pennsylvania (Received December 19, 1966)
Recently, Holroyd2 has used his 14Gmethyl radical sampling technique to look at the scavenging by ethylene of hydrogen atoms produced in the radiolysis of a number of hydrocarbons. From these studies he has (1) Supported, in part, by the U. S. Atomic Energy Commission. (2) R. A. Holroyd, J. Phys. Chem., 70, 1341 (1966).
Volume 71, Number 6 May 1967
