GEORGES CHILTZ,CARLF. ATEN,JR.,ASD S. H. BAUER
1426
TABLE I11 ACTIVITYCOEFFICIENTS O F HBr AT 0, 25, 7----y= ?"
a
U
--._
(on)----
b
-.
ASD
50"; L?AT 25'
y = (26')----
a
Vol. 66
-Lz
(ZOO)---
b
a
b
(25') cal. mole-]b
0,005 0.932 0.932 0,9295 0.9295 0.926 0.926 44 47 ,01 ,910 ,910 I906 ,90G ,902 ,602 61 64 ,03 ,884 ,883 ,879 ,879 ,8735 873 81 85 ,05 ,845 ,843 ,838 ,838 ,831 829 118 124 .07 ,832 .. ,823 ... ,8155 .. 136 .. .10 ,8165 ,812 ,808 ,805 ,797 794 168 163 This investigation. b Data of Harned, Keston, and Donelson.3 Values a t 50" were interpolated from data at 45 and G O o ,
report'ed in this paper. Again, the only explanation we can advance for the difference is an unsuspected difference in t'he method of preparing the electrodes. It has been proposed13 that the silver-silver chloride electrode be standardized by measuring(13) R, G. Bates, E. A. Guggenheim, H. s. Harned, D. J. G. Ives, G . J. Janz, C. B. Monk, J. E. Prue, R. $. Robinson, R. H . Stokes, and W. F. K. Wynne-Jones, J. Chem. Phge., 2 6 , 361 (1956).
the potential of the cell: Hz; 0.01 m HC1, AgC1; Ag a t 25' and using ~ ~ ( 0 . 0m1 HC1) = 0.904. Although the standard potentials of the silversilver bromide electrode of H, K, and D and those now reported do not agree with t'hose of Harned and Donelson, all three sets of measurements lead to rk(O.Ol& HBr) = 0.906 at 25'; thus the silverbromide can be ized in the 0.01 m solution of hydrobromic acid.
HATE OF DECO;C1POSITIBN OF AZOMETHhNE I N h SHOCK TUBE B Y GEORGES CI-IILTZ,~ CARLF. A . T E K , JR.,S N D S. H. BAUER Department of Chemistry and the Graduate School of Aeronautzcal Engzneerzny, Cornell Unzverszty, Ithaca, S e w Yorlz Recewed February 16, 1962
The pyrolysis of azomethane was studied in a shock tube over the temperature range 800-1300°K. Concentrations of the asomethane of 1-370 in argon were used. The rate of decomposition in the incident ahock region mas followed spectropiiotometrically a t 338 mp. It was estimated that under shock conditions the depletion of the reactant b z a a chain mechan,sm was negligible compared to that due t o the unimolecular decomposition. Due to the exothermicity of the over-all reaction (products, Nz and CzHC), only average rate constants could be evaluated from the recorded oscilloscope traces. These were found to fall well within the extrapolations of the two most recent low-temperature studies, based on a strict Arrhenius temperature dependence.
Introduction The thermal decomposition of azomethane ofteu is cit,ed as a classical example of a gaseous unimolecular reactmion. In 1927, Ramsperger reported2 an activation energy of 31.2 kcal./mole for this decompositioll. His lat8er experimental results gave a rate constant IC" = 3.13 X 10l6 exp(-52,500/RT). A value close to this one is still a ~ c e p t ? d . ~ Yo - ~ one has succeeded in studying this decomposition at' a sufficiently low pressure to detect' the t'ransition to the collisionlimit2d regime. Recently, however, this decompositio:i has b-en added to the already large list of "unimolecular" processes which proved to be more complex than indicated. Steel and Trotman-Dickensone found for t'he rate constant IC" = 1015.7 exp( -51,20O/RT) but emphasized t'hat a chain reaction occurs simultaneously with the uiiimolecular process. At 56OOK. and 60 mm. (1) Meurice Institute of Technology, Brussels. (2) H. C. Ramsperger, J . A m . Chem. Sac., 49, 912, 1495 (1927); SO, 123 (1928).
(3) (a) 0. K. Rice and D. V. Sickman, J . Chem. PhUs., 4, 239, 242, 508 (1936); (b) E. W. Ribbett a n d C. Rubin, J . A m . Chem. Soc., 69, 1537 (1937). 14) 131. Page, H. 0. Pritchard, a n d A. F. Trotman-Diekenson, J . Chem. Soc.. 3878 (1953). ( 5 ) M. H. Jones and E. W. R. Steacie, J . Chem. Phys.. 21, 1018 (1953). (6) C. Steel and A. F. Trotman-Diokenson,
(1959).
J. Chem.
Soc., 975
pressure, the rate of decomposition due to the radical chain is approximately equal to the unimolecular rate. That a chain is involved is inherently plausible. Since the initial products of the decomposition are methyl radicals, these not only can associat-: with each other but also may attack the residual azometharie. This secondary reaction of azomethane can be suppressed by the addition of radical inhibitors, such as nitric oxide, propylene, and toluene. Such an experiment has been reported by Forst and Rice.' They found log IC" = 17.3 - 55,400/(2.303RT). X comparison of the two most recent studies is given in Fig. 1. The temperature range covered in the abore investigations is 500-600°K. Due to ths rather high preexponential factor, it has been proposed that the fragmentation step is a simultaqeous fission of the two K-CHS bonds. The AHoo expected for such a process is approximately 15.7 kcal. as compared with the observed activation energy of around 53 kcal. and an overall AHoo of -67.6 kcal. for the productioii of ethane and nitrogen. It is of interest to study the thermal decornpositioii of azomethaiie a t temperatures above 6OO0K. for several reasons. Although the detailed mechanism for the radical chain has not (7) W. Forst and 0 K. Rlce, I b s t r . No. 118, DIV. Phys Chem , 139th National Meeting, ACB, St. Louis, N o . , ;\larch, 1061.
August, 1982
RATEOF DECOMPOSITION OF AZOMETHANE IN A SHOCK TUBE
been established, one can readily formulate a simple sequence of reactions which provides a framework for semi-quantitative estimates. One result of such an analysis is that a t 800OK. the extent of decomposition of the azomethane due to the radical attack is expected to be only 1% of the uiiimolecular decomposition. This follows from the lower activation energy for the chain reaction, which was assumed t o be 35 kcal. as compared with the considerably higher value for the unimolecular process. Secondly, extension of the observed temperature range for the unimolecular step hopefully would provide a more precise value for the activation energy. The most crucial objective is the measurement of the rate constant over a large range of temperatures to establish the validity of a strict Arrhenius temperature dependence. I n turn, this would provide a measure of the relative efficiency for decomposition as a function of the energy which a molecule possesses above the lower limit required for dissociation. It has been established8 that a temperature independent preexponential factor implies that the probability for dissociation is II-E
=
if E
0
