T H E THERMAL DECOMPOSITION OF NITROMETHANE' H. AUSTIN TAYLOR AND VLADIMIR V. VESSELOVSKY Nichols Chemical Laboratory, New York University, New York C i t y Received November 84, 1934
In attempts during recent years to isolate data for unimolecular reactions the decomposition of various organic halides, ethers, azo compounds, amines, and nitrites has been studied kinetically, and though superficial results would suggest that in most cases a net reaction of the first order was occurring, extremely few are so free from secondary changes that the results may be accepted unequivocally. The multiplication of such attempts is thus justified. No nitro compounds have so far been studied. The simplest, nitromethane, was therefore chosen. The apparatus and method used for the study were identical with those used previously by Taylor (3) in similar work, and involved the determination of the rate of pressure change of the reactant with time. To prevent condensation of the nitromethane vapor the apparatus outside the furnace was maintained a t about 80°C. throughout the work. The nitromethane used was carefully fractionated from a Kahlbaum sample, the fraction boiling between 100.5 and 101°C. being collected. Temperatures from 390 to 420°C. were found to yield a convenient velocity of decomposition. The percentage increase in pressure during reaction was found to be 130, independent of temperature and pressure, as is shown in table 1. Data of a typical experiment are given in table 2 showing the observed pressure increases occurring at the specified times a t 420°C. with an initial pressure of 198 mm. of nitromethane. The complete data are presented in table 3 in the form of fractional lives calculated as the times necessary for 25, 50, and 75 per cent of the total pressure increase to occur. This procedure is justified, since from the constancy of the end points under all conditions studied the same reaction is proceeding in each case. From the general constancy of the above values, particularly for the quarter-lives a t the higher temperatures, the reaction may be taken as of the first order, a t least early in the reaction. The changing values of the three-quarter-lives would suggest the presence of secondary reactions of higher order than the first. I n view of this possibility, the observed inAbstract from a thesis presented in partial fulfillment of the requirements for the degree of Doctor of Philosophy a t New York University, June, 1934. 1095
1096
H. AUSTIN TAYLOR AND VLADIMIR V. VESSELOVSKY
crease in the quarter-lives with decrease in initial pressure especially at the lower temperatures may be due either to a real deviation of the reaction TABLE 1 Increase in pressure during decomposition of nitromethtne INITIAL PRESSURE
I
PER CENT INCREASE
11
I
INITIAL PRESSURE
PER CENT INCREME
Temperature = 390°C. 201 148 100.5
129 130 130
55.5 25.5
131 133
Temperature = 400°C. 201.5 156.1 101.5
130
1
52 26.5
Temperature = 410OC. 187.5 145 105.5
130
(I
:,5
Temperature = 420°C. 198 157 104
130
51
'
1 1
1
133 130
130 131
130 129
TABLE 2 Data of a typical experiment Temperature, 420°C.; initial pressure, 198 mm. PRESSURE INCREASE
TIME
PRESSURE INCREASE
TIME
minutes
mm.
minufes
mm.
25 30 35 40 45 75 334 469 1354
234.5 240 243.5 246 247 249.5 251.5 253.5 257
1 2 3 4 5 6 8 10 15 20
40 67 89 108 124 138 161 179 208 224.5
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THERMAL DECOMPOSITION OF NITROMETHANE
reaction, would have a relatively larger effect on the overall rate at lower temperatures and higher pressures. TABLE 3
Complete data for the decomposition expressed as fractional lives INITIAL PRESSURE
mm.
198 157 104 51 27
1
FRACTIONAL LIVES 126
I
tm
I
176
minutes
minules
minutes
1.40 1.40 1.50 1.40 1.38
4.00 4.00 4.00 3.50 4.00
7.50 7.00 7.50 7.50 7.50
Temperature = 410°C. 7.40 7.40 7.80 7.60 7.40
189 153 97.5 46.5 23.5
2.50 2.60 2.60 2.70 2.60
15.4 15.2 15.4 15.6 15.4
201.5 156 101.5 52 26.5
5.10 5.12 5.50 5.20 5.30
14.0 14.0 14.0 13.0 14.0
28.0 28.0 28.0 25.0 24.0
307 202.5 148 100.5 55 25.5
6.50 8.00 10.40 10.35 10.40 10.40
21 .o 27.0 27.5 29.0 32.0
42.0 51.6 52.1 61.0 63.0
199.5 146 101 49.5 24.5
16.0 16.5 17.5 21.0 22.5
1098
H. AUSTIN TAYLOR AND VLADIMIR V. VESSELOVSKY
To find the effect of increased surface the reaction vessel was filled with short lengths of Pyrex tubing, the ratio of surface to volume being increased 6.3 times. A comparison between experiments on the normal and increased surface is afforded by a typical run as shown in table 4. 'Taking the total pressure increase as 130 per cent for the reaction in the packed vessel as in the empty vessel, the quarter- and half-lives €or the above reactions are 1.40 and 4.00 minutes in the enipty vessel and 1.45 and 4.10 minutes in the packed vessel. The reaction in its early stages does not therefore appear to be influenced to any marked extent by surface. TABLE 4 Comparison between experiments on the normal and increased surface Temperature, 420'C. INITIAL PRESSURE TIME
Empty, 198 mm.
1
Packed, 216.5 mm.
Prassure increase minutes
mm.
1 2 3 4 5 6 8 10
40 67 89 108 124 138 161 179 224.5 246 247 248.5 249.5 250 25 1 257
20 40 50 . 60
80 170 260 1350
?am2.
50 87.5 115 136 152 162 181 190 205 206 202 199 193 188 192 199
That the later reactions, however, are changed by the extent of surface can readily be seen from the peculiar pressure changes in the packed vessel given above, wherein the pressure increase reaches a maximum, falls, and then later rises again. This pressure decrease always observed in the packed vessel, though never found with the empty vessel, can be traced, as is shown later, to a condensation in the capillaries of some of the products of reaction. This fact is responsible for the apparent changed end point in the packed vessel. The influence of foreign gases in the course of reaction was very thoroughly investigated for the following gases,-helium, nitrogen, nitric oxide, carbon dioxide, and oxygen. With the exception of oxygen all
1099
THERMAL DECOhiPOSITION OF NITROMETHANE
these gases were found to be without effect, the observed results, indeed, duplicating those already given. The addition of oxygen resulted in an appreciable change of the reaction rate and a somewhat different end point. The data are presented in table 5 . The energy of activation of the reaction in its early stages was found by application of the Arrhenius equation to the average values of the quarterlives at the higher pressures as previously given. The logarithms of these times plotted against the reciprocals of the absolute temperatures gave a good straight line with a slope corresponding to 61,000 calories. The actual mechanism of the reaction seems to be so extremely complex that it will be simplest to outline it step by step and to attempt to justify 'each step as given. The primary step postulated involves a split of oxygen according to the equation CHBNOz + CH,NO
TIME
PRESSURI INCREASE
minutes
mm.
1 2 3 4 5
29.5 49.5 65 77.5 87.5 99.5 106 113.5
6 8 10
+
+02
TIME
~
I
PRESSURE INCREASE
minules
mm.
12 15 20 40 70 250 1350
118.5 122.5 125.5 126 126 129 130
That nitrosomethane is formed fairly readily was demonstrated by refluxing nitromethane a t its boiling point for forty-eight hours. The liquid was then fractionated and a small fraction boiling a t 84°C. collected. This is the boiling point of formaldoxine, the isomer of nitrosomethane. Upon refluxing a small portion of this liquid with water for some hours and subsequent addition of silver nitrate a copious precipitate of silver cyanide was obtained, according to the reaction CHz:NOH
-+
HCN
+ HzO
Another portion of the formaldoxime was hydrolyzed in presence of acid and gave a subsequent test for aldehyde CHz:NOH
+ HzO
-+
CHzO
+ NHZOH
1100
H. AUSTIN TAYLOR AND VLADIMIR V. VESSELOVSKY
The proven presence of formaldoxime is indirect substantiation for the postulated nitrosomethane which is known to be unstable and to isomerize as stated (2). Nitromethane is known to oxidize according to the reaction (l), CH3NOz
+ ~ O-+Z GO2 + $HzO + $Nz
and gas analyses of the products of reaction showed their presence in large quantities. The reaction indeed must b e a rapid one, as is demonstrated by the increased rate observed in the experiments with added oxygen as compared with those in its absence. The third step involves the fate of the nitrosomethane; the simplest assumption would be the splitting into nitric oxide and free methyl radicals,. which would give ethane alone or a mixture of methane, ethylene, and hydrogen, as frequently appears in hydrocarbon decompositions. Assuming the former, 2CH3NO 4 CzHa
+ 2NO
there would be for the overall reaction occurring 10CH3N023 6 N 0
+ 6H20 + 4C02 + 3CzHa + 2Nz
The presence of large amounts of nitric oxide continually made itself felt during the conduct of the experiments. Copious brown fumes of the dioxide were always observed upon evacuating the system after each run. In an effort to identify the above products and their relative amounts a t the end of the decomposition a number of bombs were made up containing nitromethane; they were then heated to allow complete reaction to occur. The bombs upon cooling always showed large amounts of water. In some, definite tests for cyanides were obtained and others contained a white solid. A microanalysis of the solid2 gave it the empirical formula CH6N03,and it later proved to be ammonium bicarbonate. Analysis of the gases remaining in the bomb gave on the average 23 per cent carbon dioxide, 30 per cent carbon monoxide, 16 per cent methane, 4 per cent nitric oxide, about 1 per cent each of hydrogen and an unsaturated hydrocarbon reckoned as ethylene, with the residual 25 per cent of nitrogen. These analyses, bearing only superficial resemblance to the amount of products postulated above, suggested the possibility that concentration and surface conditions existing in the bombs might have a marked effect on the later progress of reaction. It will be recalled that although no effect of increased surface was found early in the reaction, the later stages were greatly changed. An observation of significance too, was that instead of copious brown fumes being observed on evacuating the system after runs
* Thanks are due Dr. Joseph
B. Niederl for this analysis.
THERMAL DECOMPOSITION OF NITROMETHANE
1101
with'the increased surface, vapors with a strong ammoniacal odor alone were found. Close inspection of the capillaries after a number of these runs showed traces of the white solid, ammonium bicarbonate. Apparently one of the later reactions in the series is capable of being catalyzed to yield varying products; it is suggested as most likely that the nitrosomethane decomposition is responsible, since it is this reaction which yields the observed nitric oxide in large amounts with the empty vessel. The condensation of ammonium bicarbonate in the capillaries accounts satisfactorily for the pressure decrease observed with the increased surface and the reduced end point as compared with runs in the empty vessel. As stated previously, nitrosomethane will isomerize to formaldoxime, which in the presence of water from the nitromethane oxidation will hydrolyze to yield hydroxylamine and formaldehyde. The latter would be decomposed at the temperature in question into carbon monoxide and hydrogen. The hydroxylamine would yield ammonia, and hence give rise to ammonium bicarbonate in the system. It seems quite reasonable then, that the specific surface effect may be in the isomerization of the nitrosomethane to formaldoxime and, in view of the large amounts of nitric oxide produced in the runs in the unpacked vessel, to assume the mechanism given above. The overall reaction would then correspond to a pressure increase of 110 per cent as compared with the observed 330 per cent. If the reactions of the free radicals from the nitrosomethane should not give ethane alone, and the 16 per cent of methane in the bomb experiments would seem to suggest this, the total pressure change would be greater than that given above and more nearly in agreement with the observed. In summary then, the main reaction being studied appears to be a homogeneous unimolecular split of nitromethane into nitrosomethane and oxygen, with an energy of activation of about 61,000 calories. The complexity of the subsequent reactions does not permit a further analysis of the observed kinetic data. REFERENCES (1) BERTHOLET: Ann. chim. phys. [6] 30, 567 (1893). (2) SCHOLL: Ber. 24, 576 (1891). J. Phys. Chem. 34, 2761 (1930). (3) TAYLOR:
