THE PRODUCTION OF THE LOWER CHLORIDES OF METHANE FROM NATURAL GAS
Journal of Industrial & Engineering Chemistry
The large amount of methane available in the form of natural gas has made t h e transformation of the same into chloroform an attractive problem. The low cost of carbon tetrachloride from other sources, along with the fact t h a t in t h e chlorination of methane half of the halogen goes t o form hydrogen chloride, has made the formation of the tetrachloride from methane less interesting. Baskerville and Hamor2 have very completely covered the literature on this subject, as well as the
... ell as the manufacture of chloroform from materials other t h a n methane. The following references, however, are repeated, since they are closely related t o this article. B e r t h e l~t ,~ in 18j8, showed t h a t unless t h e reaction between chlorine and methane was carried out with extreme slowness, explosions and the separation of carbon easily resulted. P h i l l i p~,~ in 1893, tried t o avoid explosions b y chlorinating methane, without the access of light, in a tube heated t o 3o0-400~ C., and finding t h a t only the first and last chlorides of the series were formed in appreciable amounts, he ques-tioned5 the successful chlorination t o intermediate products. Tolloczko,G in 1912, also working in hot tubes, obtained results similar t o Phillips. Walter,' in 1909, states t h a t water vapor cuts down the speed of chlorination considerably, and if the chlorine is added gradually t h a t it reacts more readily with the partially chlorinated methane than with the original hydrocarbon, the result being the highest chlorine substitution possible. Baskerville and Riederers cite t h e work of Phillips and Walter and seek t o obtain chlorofbrm indirectly by chlorinating methane t o the tetrachloride and then reducing. They also show t h a t the blue end of the visible spectrum is more suitable for use in chlorination t h a n ultraviolet rays. G r a d and Hanschke,gin 1912, state t h a t the chlorination of ethane in diffused daylight, as carried out by SchorIemmer,'o in 1869, is impractical, since a t low temperatures the halogen reacts very slowly. Thermochemical calculationsll predict a rise of over 2oooo C. for the reaction of a mixture of methane and chlorine in the proportions t o form carbon tetrachloride, provided all the heat of reaction goes t o increase the temperature of t h e reaction products; 1 Patent applications have been made, covering the work outlined in this article. 11 Using 21,000 as the heat in the formation of carbon tetrachloride. (See "Thermochemistry," Thomsen-Ramsay series, p. 246.) Berthelot and Matignon give 68,500 as the heat of formation of carbon tetrachloride, which gives a value of over ,3000" C. in the above calculation for equal volumes of methane and chlorine, the rise calculated for CH3Cl is over 1700'; for a mixture of j vols. methane : I vol. chlorine, .the theoretical rise is nearly 600' so t h a t even with this high dilution of the gases, if the reaction were carried out in a tube already heated to 400 O , the resulting temperature would approximate I O O O O C. At or above 1000' C.. methane and its derivatives will decompose, and Phillips st,ates t h a t he frequently observed the deposition of carbori inside his heated tube. A mixture of 4 vols. of chlorine with I vol. methane, the theoretical proportions t o form carbon tetrachloride, when exposed t o direct sunlight in glass tubes of small diameter, will not give a violent reaction and if the air is cold. the reaction will proceed slowly. The gases are so closely in contact with the tube walls t h a t any heat liberated is radiated a t once and the temperature cannot rise t o t h a t of speedy react:on even in direct sunlight. A mixture of chlorine and methane, in the same proportions in a large balloon flask and in direct sunlight, will explode with great violence. The interior of the gas body, being heatinsulated by the surrounding layers of gas, quickly comes u p t o the temperature of speedy reaction and then rises suddenly t o I O O O O C. or' higher, where carbonization begins t o take place. At temperatures even as low as 100' C., all of the chlorine derivatives of methane will remain in the vapor state and with chlorine showing a preference for the partially chlorinated hydrocarbon, t h e intermediate compounds cannot easily be obtained. The problem, therefore, is t o obtain a speedy chlorination of methane while maintaining an average temperature low enough t o condense out the chloroform before it can be chlorinated further. At high temperatures no light is necessary t o induce t h e reaction between chlorine and methane. In order to maintain a low temperature, the amount of free chlorine in the gases a t any time must be correspondingly low, and the lower the content of the chlorine, the stronger the light necessary t o induce a speedy reaction a t this low temperature. I t therefore became evident, from a few preliminary experiments, t h a t the successful solution of the problem was partly dcpendent on the use of a sufficiently powerful light. The "White Flame Arc," as described by Mott and Bedford,' was found t o be the most satisfactory source of light and was used in all of the following work.