Pyrolysis of methyl chloride, a pathway in the chlorine‐catalyzed polymerization of methane

The reaction of CH 4 + Cl 2 produces predominantly CH 3 Cl + 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 CH 3 Cl formed C 2 H 4 and C 2 H 2 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 CH 4 C 2 H 4 and C 2 H 2 are formed by the recombination of ĊH 3 and ĊH 2 Cl radicals. With added CH 4 recombination of ĊH 3 forms C 2 H 6 , which dehydrogenates to C 2 H 4 + H 2 . C 2 H 4 in turn dehydrogenates to C 2 H 2 + H 2 . While HCl, C, CH 4 , and H 2 are the ultimate stable products, C 2 H 4 , C 2 H 2 , and C 6 H 6 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. ĊH 3 ‐ initiated polymerization of ethylene is negligible relative to the Ċ 2 H 3 formation through H abstraction by Cl. The fastest reaction of Ċ 2 H 3 is its decomposition to C 2 H 2 . About 20% of the consumption of C 2 H 2 can be accounted for by the addition of Ċ 2 H 3 to it with formation of the butadienyl radical. The addition of the latter to C 2 H 2 is slow relative to its decomposition to vinylacetylene. Successive H abstraction by Cl from C 4 H 4 leading to diacetylene has rates compatible with the experimental values. About 10% of Ċ 4 H 5 abstracts H from HCl and forms butadiene. Successive additions of Ċ 2 H 3 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 Ċ 2 H 3 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}