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Notes: It is shown that chain reactions may be studied in stirred-flow reactors under conditions in which the escape of radicals from the reactor competes significantly with the chemical termination reaction. The appropriate rate law permits the independent evaluation of the rate constants of the initiation and termination steps. Such information is of special interest in the study of surface effects on chain reactions.

Notes: NO Abstract.

Notes: Using published data on the kinetics of pyrolysis of C2Cl6 and estimated rate parameters for all the involved radical reactions, a mechanism is proposed which accounts quantitatively for all the observations:The steady-state rate law valid for after about 0.1% reaction is and the reaction is verified to proceed through the two parallel stages suggested earlier whose net reaction isA reported induction period obtained from pressure measurements used to follow the rate is shown to be compatible with the endothermicity of reaction A, giving rise to a self-cooling of the gaseous mixture and thus an overall pressure decrease.From the analysis, the bond dissociation energy DH0(C2Cl5—Cl) is found to be 70.3 ± 1 kcal/mol and ΔHf3000(·C2Cl5) = 7.7 ± 1 kcal/mol. The resulting π—bond energy in C2Cl4 is 52.5 ± 1 kcal/mol.

Notes: An analysis of the thermochemistry of the kinetic parameters of the elementary reactions involved in the pyrolysis of pentachloroethane has resolved several disputed, unclarified, or inconsistent aspects of the reaction mechanism. The resulting mechanisms for the inhibited and uninhibited pyrolysis account for all reported experimental findings. On the basis of this interpretation, first experimentally based values have been derived for the following: DH0(CCl3-CHCl2) = 79.0 ± 1.0 kcal/mol, ΔHf0(CHCl2) = 25.7 ± 1.0 kcal/mol, and E1 = 59.7 ± 1.0 kcal/mol C2HCl5 .

Notes: The reaction of CH4 + Cl2 produces predominantly CH3Cl + HCl, which above 1200 K goes to olefins, aromatics, and HCl. Results obtained in laboratory experiments and detailed modeling of the chlorine-catalyzed polymerization of methane at 1260 and 1310 K are presented. The reaction can be separated into two stages, the chlorination of methane and pyrolysis of methylchloride. The pyrolysis of CH3Cl formed C2H4 and C2H2 in increasing yields as the degree of conversion decreased and the excess of methane increased. Changes of temperature, pressure, or additions of HCl had little effect. In the absence of CH4 C2H4 and C2H2 are formed by the recombination of ĊH3 and ĊH2Cl radicals. With added CH4 recombination of ĊH3 forms C2H6, which dehydrogenates to C2H4 + H2. C2H4 in turn dehydrogenates to C2H2 + H2. While HCl, C, CH4, and H2 are the ultimate stable products, C2H4, C2H2, and C6H6 are produced as intermediates and appear to approach stationary concentrations in the system. Their secondary reactions can be described by radical reactions, which can lead to soot formation. ĊH3 - initiated polymerization of ethylene is negligible relative to the Ċ2H3 formation through H abstraction by Cl. The fastest reaction of Ċ2H3 is its decomposition to C2H2. About 20% of the consumption of C2H2 can be accounted for by the addition of Ċ2H3 to it with formation of the butadienyl radical. The addition of the latter to C2H2 is slow relative to its decomposition to vinylacetylene. Successive H abstraction by Cl from C4H4 leading to diacetylene has rates compatible with the experimental values. About 10% of Ċ4H5 abstracts H from HCl and forms butadiene. Successive additions of Ċ2H3 to butadiene and the products of addition can account for the formation of benzene, styrene, naphthalene, and higher polyaromatics. The following rate parameters have been derived on the basis of the experimentally measured reaction rates, the estimated frequency factors, and the currently available heat of formation of the Ċ2H3 radical (69 kcal/mol): \documentclass{article}\pagestyle{empty}\begin{document}$$ \begin{array}{*{20}c} {\mathop {{\rm C}_{\rm 2} }\limits^. {\rm H}_{\rm 3} \mathop {\longrightarrow}\limits_{\left( {\rm M} \right)}^{39} {\rm H}\,\, + \,\,{\rm C}_{\rm 2} {\rm H}_{\rm 2} } & {\log k\left( {1\,{\rm atm,}\,{\rm 1300}\,{\rm K}} \right)\, = \,5.2\, + \,0.3\,s^{ - 1} } \\ \end{array} $$\end{document} \documentclass{article}\pagestyle{empty}\begin{document}$$ \begin{array}{*{20}c} {{\rm C}_{\rm 2} {\rm H}_{\rm 4} \, + \,\mathop {{\rm C}_{\rm 2} }\limits^. \,\mathop {\longrightarrow}\limits^{17} \,\mathop {{\rm C}_{\rm 4} }\limits^. {\rm H}_{\rm 7} } \hfill & {E\, \ge \,2\, \pm \,2\,{{{\rm kcal}} \mathord{\left/ {\vphantom {{{\rm kcal}} {{\rm mol}}}} \right. \kern-\nulldelimiterspace} {{\rm mol}}}\,} \hfill \\ {\mathop {{\rm C}_{\rm 2} }\limits^. {\rm H}_{\rm 5} \, + \,{\rm C}_{\rm 6} {\rm H}_{\rm 6} \,\mathop {\longrightarrow}\limits^{40} \,\mathop {{\rm C}_{{\rm 12}} }\limits^. {\rm H}_{{\rm 11}} } \hfill & {E\, = \,11\, \pm \,2\,{{{\rm kcal}} \mathord{\left/ {\vphantom {{{\rm kcal}} {{\rm mol}}}} \right. \kern-\nulldelimiterspace} {{\rm mol}}}} \hfill \\ \end{array} $$\end{document}