might equal 0.05 ppm. If it is assumed that NO and NO2 are equal in concentration and that the amount of water vapor present is 13 000 ppm, the concentration of nitrous acid would be 0.0035 ppm in the dark. Any trace of dimethylamine introduced into such an atmosphere would react a t a rate of about 2%/h. For heavier urban pollution, the reaction rate would be correspondingly higher. The above calculations apply only in the dark. When the sun rises, the nitrous acid is destroyed by photolysis much faster than it can be reformed by the relatively slow termolecular reaction of NO, NOz, and H 2 0 . Nitrosation therefore cannot take place in the daylight. The same light that stops the nitrosation process will destroy the accumulated N-nitroso dimethylamine. As reported above, at full sunlight half the nitrosamine will be destroyed in IJ2h. Between 8:OO a.m. and noon, any N-nitroso dimethylamine accumulated during the night would be reduced to less than one-half of 1%of its starting concentration. After 10 h of daylight, only one molecule in one million will remain. The study of photolysis products is not complete. To date, the results have shown that the destroyed N-nitroso dimethylamine is replaced by a mixture of nitric oxide, carbon monoxide, formaldehyde, and an unidentified compound whose spectrum is shown in Figure 3. A weakening of the C-H band indicated that this latter compound may only contain a small fraction of the starting material. The spectrum of products does not exhibit any 0-H stretching frequency, thus indicating that a molecular rearrangement to an oxime does not take place. In this respect, the gas-phase photolysis diffefs from the liquid-phase photolysis reported by Chow (5). The large nitric oxide yield indicates that the photolysis breaks the N-N bond: (CH:j)pN-NO
+ h~
+
+
(CH:))?N NO
The radical (CH:j)pN most likely gives up methyl radicals that are further oxidized to the CO and H&O seen as products. It seems possible that the unidentified compound is the result of oxidation and/or recombination of such radicals as (CH:J2N,CH2N, and CH:).
Conclusions It seems reasonable to conclude that atmospheric formation of N-nitroso dimethylamine should not be regarded as a serious general problem. This conclusion is based on two factors. First, the reaction between dimethylamine and nitrous acid will only occur to an appreciable extent while the amine is confined to a region of heavy urban air pollution. The second factor to be considered is that dimethylamine is not a generally distributed air pollutant; the problem will arise only in the vicinity of manufacturing plants that release the amine. The indicated nonformation of the nitrosamine during the daylight hours requires that any nitrosamine detected in the air in the afternoon must have emitted as nitrosamine rather than in the form of precursors. The destruction of N-nitroso dimethylamine by sunlight will prevent day-to-dayaccumulation of the compound in the air. It seems reasonable to assume that any N-nitroso dimethylamine that may be present in particles in the air will also be destroyed by photolysis. Formation of N-nitroso dimethylamine would take place more readily in industrial atmospheres than in the ambient air, because indoor illumination does not have sufficient ultraviolet intensity for photolytic destruction of either the nitrous acid reactant or the nitrosamine product.
Literature Cited
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(1) Lucas. H. J.. “Organic Chemistrv”. American Book Co.. New York, N.Y., 1935. (2) Pauchert, C. J.. “The Aldrich Library of Infrared Soectra”, 2nd ed., Aldrich Chemical Co., Milwaukee; Wis., 1975. ( 3 ) Glasson, W. A., General Motors Research Lab, Warren, Mich., oublic comments made at the 171st Meeting. ” ACS. New York. N.Y.. April 4-9, 1976. 14) Calvert. J. G.. Chan. W. H., Nordstrom. R. J., Shaw, J. H., “Spectroscopic Studies of Photochemical Smog and Trace Pollutant Detection”, Ohio State University, Columbus, Ohio, 1975. (5) Chow, Y.-L., Tetrahedron Lett., 33-34,2333 (1964). I
,
Received /or reuieu July 6, 1976. Accepted October 15, 1976. Presented at the Division of Environmental Chemistry, 171st Meeting, ACS, N e u York, N . Y., April 4-9, 1976.
Interphase Transfer Processes. II. Evaporation Rates of Chloro Methanes, Ethanes, Ethylenes, Propanes, and Propylenes from Dilute Aqueous Solutions. Comparisons with Theoretical Predictions Wendell L. Dilling Environmental Sciences Research Laboratory, The Dow Chemical Co., Midland, Mich. 48640
Two theoretical models for predicting the rate of evaporation of slightly soluble materials from aqueous solution have been published recently (1-3). In a continuation of our work on the evaporation of low-molecular-weight chlorohydrocarbons from dilute aqueous solution [Part I ( 4 ) ] ,we have measured the evaporation rates of 27 chlorohydrocarbons from water to test these two theoretical models. Preliminary work ( 4 ) had indicated that the earlier model ( 1 ) was inadequate since neither absolute nor relative predicted rates were in agreement with the experimental rates. In the present work the evaporation rates of all of the chloro methanes, ethanes, and ethylenes, and a few of the chloro propanes and propylenes were measured to carry out a more extensive test.
Experimental The same hollow fiber-mass spectrometric procedure as reported previously ( 4 )was used to determine the evaporation rates. A schematic diagram which shows the geometry and approximate degree of turbulence of the solution is shown in Figure 1. Two to five compounds were run simultaneously in the same solution. The initial concentration of each component was 1.0 ppm (weight basis). Typical data are shown in Figures 2 and 3. The ion-peak height was correlated with concentration by extrapolation of the decay portion of the curves in Figures 2 and 3 to zero time. This extrapolated concentration a t zero time was taken as 1.0 ppm. Successive Volume 11, Number 4, April 1977
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rn The evaporation rates of 27 of the title compounds from dilute (-1 ppm) aqueous solutions were determined, and the data were used to test two theoretical models. Most of the rates followed first order kinetics for the first 2-5 half-lives. Under the experimental conditions [200 rpm stirring (shallow pitch propeller), -25 OC, still air (go% liquid-phase resistance), there is relatively little change in K I with major changes in H . For example, the forty-fivefold variation in the two values of H for CHl=CHCl in Table I leads to only a 1%change in K I and t i Thus, for compounds with H > 0.1, i.e., most but not all of the chlorohydrocarbons in this study, the calculated half-life, t i ?, is almost directly proportional to t he square root of the molecular weight and is nearly independent of H . The earlier theoretical model by Mackay and Wolkoff ( I ) leads to Equation 5 where t ,”? is the solute evaporation
’.
-
0.01797 GP,,S (5) EPM half-life in minutes, G is the weight of water in grams, P,, is the partial vapor pressure of water in mm of Hg, and E is the weight of water in grams which evaporates per day. In our experiments a t -25O, G = 200, P,, = 23.76, and E = 20. Calculated values of t ;’.are listed in Table I. Comparison of the values of t ?’: with the experimental values, ti ?, shows that Equation 5 fails to predict both the absolute and relative half-lives for evaporation. The possible reasons for this failure of Equation 5 have been discussed previously (2, 4 ) . t;‘, =
(1) Mackay, D., Wolkoff, A. W., Enuiron. Sei. Technol., 7, 611 (1973). ( 2 ) Mackay, D., Leinonen, P. J., ibid., 9, 1178 (1975). (3) Liss,P. S., Slater. P. G., Nature, 247,181 (1974). ( 4 ) Dilling, W. L., Tefertiller, N. B., Kallos, G. J., Enuiron. Sei. Technol., 9,833 (1975); ibid.. 10,1275 (1976). (5) Glew, D. N., Moelwyn-Hughes, E. A,, Discuss. Faraday Soc., 15, 150 (1953). (6) Swain, C. G., Thornton, E. R., J . Am. Chem. Soc., 84, 822 (1962). ( 7 ) McConnell, G., Ferguson, D. M., Pearson, C. R., Endeauour, 34, 13 (1975). (8) Pearson. C. R., McConnell, G., Proc. Roy. Soc. London, B189,305 (1975). (9) Hardie, D.W.F., “Kirk-Othmer Encyclopedia of Chemical Technology”, A. Standen, Ed., Vol 5, 2nd ed., pp 100-203, Interscience, New York, N.Y., 1964. (10) Syverud, A. N., The Dow Chemical Co., Midland, Mich., private communications. 1969, 1970. (11) Brown. E. J.. The Dow Chemical Co.. Plaauemine. La.. orivate communication, 1973. (12) McGovern, E. W., Ind Ena Chem., 35,1230 (1943). ( 1 3 ) McDonald, R. A,, The Dow Chemical Co., Midland. Mich., private communication, 1970. (13) Marsden, C., Mann, S.,“Solvents Guide”, pp 261, 269,507,533, 535, 540, Cleaver-Hume Press, London, England, 1963. (15) Saracco, G., hlarchetti, E. S., Ann. Chim. (Rome), 48, 1357 ( 19ci8). (16) Braker, W., Mossman. A. L., “Matheson Gas Data Book”, 5th ed.. p 241. Matheson Gas Products. East Rutherford, N.J., 1971. (17) Van Arkel, A. E., Vles, S. E., Rec. Trac. Chim., 55,407 (1936). (18) Stull, D. R., The Dow Chemical Co.. Midland, Mich., private communication, 1944. (19) Renfro, J. C.. The Dow Chemical Co., Freeport, Tex., private communication, 1967. 120) “Material Safety Data Sheet”, The Dow Chemical Co., Midland, Mich.?1975. (21) Ambrose. D., Sprake. C.H.S., Townsend, R.. J . Chem. Soc., Faraday Trans. 1, 69,839 (1973). (22) Mortimer, R. E.. The Dow Chemical Co., Sarnia, Ont., Canada, private communication, 1971. (23) Antropov. L. I., Pogulyai, V. E., Simonov, V. D.. Shamsutdinov, T. M., Zh. Fiz. Khim. USSR, 46,534 (1972);English transl., Russ. J . Ph.ss. Chem., 46, 311 (1972). (24) Hayduk, W.. Laudie, H.. J . Chem. Eng. Data, 19,253 (1974). (25) “Organic Chlorine Compounds”, Union Carbide Corp., 1960. (26) Haalev. D. M.. The Dow Chemical Co.. Midland. Mich.. Drivate communication, 1968. ( 2 7 ) Huston, R. F., The Don Chemical Co., Freeport, Tex., private communication. 1955. (28) Pilorz. B. H.. “Kirk-Othmer Encyclopedia of Chemical Technology”, A. Standen. Ed., Vol 5. 2nd ed., pp 205-15, Interscience, New York, N.1’ ...1964. (29) Haschke, E. M., The Dow Chemical Co.. Freeport, Tex., private communication, 1975. (30) Otopkov, P.P.. Bessarab, N. A,, Rotshtein, Ti. I., Dzhagatspanyan. R. V.. Zh. Prihi. Khim., 42, 2367 (19691; English transl., J . Appl. Chem. USSR, 42,2226 (1969). (31) Goring, C.A.I., “Advances in Pest Control Research”, R. L. Metcalf, Ed.. Vo! V, pp 47-84, Interscience, New York, N.Y., 1962. (32) Neely, W. B., Proceedings of the 1976 National Conference on Control of Hazardous Material Spills, p p 197-200, New Orleans, La., April 1976.
Received for ret‘ieu,June 1 , 1976. Accepted Notiember ,i,1976. Presented at the Division of Industrial and Engineering Chemistry, 172nd Meeting, ACS, San Francisco, Calif., August 29-September 3, 1976;Abstracts of Papers, No. INDE 8.
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