THE THERMAL DECOMPOSITION OF ISOPROPYLAMINE BY H. AUSTIN TAYLOR
Previous investigations on the pyrolysis of aliphatic amines' have shown that the series offers useful and interesting examples of homogeneous unimolecular reactions. These have been shown to possess a simplicity to be contrasted with the complexities observed in ether decompositions in that activation appears to be concerned only with a single vibrational bond, corresponding that is, to two square terms. It has been suggested that the rup ture is most probably in the C-N bond. This would infer as was suggested in the previous papers that a t the temperatures employed the initial reaction involved a direct splitting out of ammonia. The alternative mechanism suggested by Hurd and Carnahan2 involves a loss of two molecules of hydrogen resulting in a nitrile formation. The probability of such an occurrence would be considerably reduced with an iso-amine in comparison with a normal straight chain compound. Disregarding the fact that the effect of added hydrogen on the decompositions of ethylamine and propylamine has been shown not to be in agreement with the nitrile mechanism, the similarity between the kinetics of the iso-propylamine decomposition here studied and the previous results would itself infer the inaccuracy of such a mechanism. I t is unfortunate that the subsequent reactions of the unsaturated hydrocarbons formed by the removal of ammonia from the amines proceed at a rate which complicates the study of what now appears definitely to be the simple rupture of the C-N bond. The similarity of the results found for isopropylamine with those for propylamine do not warrant any extended account being given. The apparatus and procedure were identical with those already published. The temperature range used was from 490°C to 54oOC. Pressures from about I j mms. to 2 2 0 mms. were used. The pressure increase during reaction was somewhat greater than for propylamine, being on the average 130 percent. The unimolecular nature of the reaction is shown in a typical case by the data in Table I showing a constant quarter life at pressures above 150 mms. with increasing values corresponding to slower reaction a t successively lower pressures.
TABLE I Temperature 4 9 5 T Initital Preeaure
18.j mms. 187.4 I 60
2
I10 j 1
Quarter Life
Initial Pressure
70
2.25
2 0
36 17
2 4
2 .o 2 2
J. Phys. Chem., 34, 2761 (1930); 35, 2658 (1931). Soc., 52, 4151 (1930).
* J. Am. Chem.
Quarter Life
0
2
2.5
671
THERMAL DECOMPOSITIOX O F ISOPROPTLAMISE
The absence of heterogeneity is typified in Table I1 giving a comparison of the actual pressure changes obtained using an empty pyrex bulb and one partly filled with pyrex powder. . TABLE 11 Temperature 495°C Empty Initial Bulb 2 1 8 . 5 mms. Pressure AP Time 0 . 5 mins. 32 mms. I 50 2 76
3
96
4
111
Pyrex Powder 229 mms. AP 26 mms.
Initial Pressure Time 6
Ei%ty mms.
Pyrex powder 229 mms.
2I8.j
AP
AP 138 158
45
8
73
IO
I73
95 113
15
208
141 I59 I73 2oj
In Table I11 are given comparative data for the pressure changes obtained in presence of zoo mms. of added hydrogen and nitrogen.
TABLE I11 Temperature 49 j"C Initial Pressure of Amine Time in mins.
160 mms.
Hydrogen 169 mms.
163 mms.
AP
AP
AP
0 5
21
20
20
36 56 68 81
34 54 69
33
82
78
98
'03
95
Ill
117
1
I88 198
I71
IO
I22
'5
I44 161
20 25
30
I75 '85
Sitrogen
53 6i
IO
190
The obvious similarity of the course of the reaction in these cases is confirmation of the unimolecular nature of the reaction. The effect of added hydrogen at intial pressures of amine below I 50 mms. where the decomposition is no longer unimolecular is similar to its effect in the ether decompositions in its ability to maintain the reaction a t its first order rate. For example from Table I it can be seen than an initial pressure of amine of 47 mms. would possess a quarter life of 2.3 mms. when decomposing alone. It was found that at the same temperature of 495"c,47 mms. of amine in presence of 186 mms. of hydrogen decomposed 2 5 percent in 2.0 mins., the value of the quarter life at higher pressures of amine in the absence of hydrogen. The time of quarter decomposition has been used in the above considerations as the criterion of the rate of reaction for reasons similar to those mentioned in the previous works. The lack of constancy of the times of j o and
672
H. AUSTIN TAYLOR
7 j percent decomposition due to secondary reactions is reflected in the falling values obtained for the unimolecular velocity constants calculated for each case examined. The behaviour is again similar to that found with propylamine. A typical example is given in Table IV.
TABLE IV Temperature p,o°C
I
I02
k 0.517 0.434
1.5
127
0.384
IO
245 259
2
146 I73 I95
0.349 0.303 0.279
15
277
20
288
Time
AP
0.5
66
3 4
Time
AP
6
225
8
k 0.249 0.233
0.224 0.2q7 0.249
The values of the quarter lives of the reaction in the stable region of pressure a t different temperatures are given in Table V.
TABLE V Temperature Quarter Life
490
495
500
510
j
2.0
1.7
1.2
2 .
520 0.9
530 0.6
54Ooc 0 . 4 5 mins.
The logarithms of these times plotted against the reciprocals of the absolute temperatures yield a straight line with a slope corresponding to an energy of activation of 42,600 calories. This value is somewhat less than t'he 44,400 calories obtained for propylamine, which was to be expected since an extrapolation of the above data shows the isopropylamine decomposition to proceed at 48ooC a t the rate attained by propylamine at 5ooOC. At this temperature of 480°C the value of E / R T is 28.4 as compared with values of 28.2 for ethylamine and 28.9 for propylamine at comparable temperatures. The final point of similarity lies in the agreement between the calculated ratio of effective to total collisions at the pressure where the reaction ceases to be unimolecular and the value of Since the reaction deviates from its unimolecular course below 1 5 0 mms. there will be a t 480OC. approximately 2 X IO*' collisions per cc. per second. The number of molecules reacting is 4 X 1 0 ' ~giving a ratio of 2 X IO-'?. The value of is 4.6 X 10-13, yielding an agreement which could be expected only if a single bond involving two square terms was responsible in the activation process for the isopropylamine as found previously for propylamine. Further work on di- and trialkyl amines is in progress.
s-aty The pyrolysis of isopropylamine has been shown to parallel that of propylamine as a homogeneous reaction, unimolecular at pressures higher than I j o mms. and possessing an energy of activation of 42,600 calories accountable on the basis of an activation and rupture of a single vibrational bond probably between the carbon and nitrogen. Nichols Chemical Laboratory, New York C'niveraity, New York, N . Y .
