T H E THERRIAL DECORIPOSITIO?; O F PROPYLAMIKE BY H. AUSTIN TAYLOR AND HAROLD E . ACHILLES]
The decomposition of ethylamine recently studied by Taylor? was shown to be a homogeneous unimolecular reaction, no deviation from this being observed down to pressures of so mms. By comparison with corresponding work on ethers it seemed probable that a more complex amine might show deviations from unimolecularity a t higher pressures and thus afford additional data whereby the various theories proposed to account for such reactions might be tested. Propylamine was chosen as a more complex though still straight-chain amine. The pyrolysis of propylamine has been studied by Upson and Sands3 by a dynamic method, and still more recently that' of ethylamine by Hurd and Carnaham.' Such a complete analysis of the complex products of decomposition is thus presented that no further analyses have been made. I n the decomposition of ethylamine Hurd and Carnahan find hydrogen, acetonitrile, methane, ethane, hydrogen cyanide, ethylene and ammonia. Correspondingly then, one would expect from propylamine, propylene, propane, ammonia, ethyl cyanide, ethane, hydrogen and hydrogen cyanide. The mechanism proposed by these workers for ethylamine pyrolysis involves a primary breakdown of the amine with the formation of an aldimide:
+ Hz
CHaCHzNH? -+ CHsCH =E" which subsequently loses hydrogen: CHsCHzSH
-+
CH3CS
+ Hz
At higher temperatures in the neighborhood of I O O O ~ the C . aldimide may yield methane: CH3CH:SH --+ CH, HCS'
+
The presence of ethane and ammonia is supposed to be due to reduction of the amine by the hydrogen formed in the primary decomposition. The reaction is written by them as:
+
C H ~ C H ~ S H ZZH + CHiCH3
+ SHa
although there is no comment indicating whether some especially active form of hydrogen would be necessary for this reaction. Upson and Sands had suggested in the initial breakdown of the amine, the formation of ammonia and the ethylidene radical, a reasonable assumption in view of the possible dehydration of the corresponding alcohol. Failing to find, in the decomposition of benxylamine, appreciable amounts of stilbene 'Abstract from a thesis presented in partial fulfillment of the requirements for t h e degree of Doctor of Philosophy a t New York University. 2 J. Phys. Chem., 34, 2761 (1930). J. Am. Chem. Soc., 44,2306 (1922). J. Am. Chem. Soc., 52, 4151 (1930).
THE THERMAL DECOMPOSITION OF PROPYLAMINE
2659
which would be expected from the combination of two phenylmethylene radicals, Hurd and Carnahan see no necessity for the postulation of the ethylidene radical in ethylamine. The similarity of the decomposition products from both alkyl and aryl amines is alone judged as evidence of a similar decomposition mechanism. According to the Hurd mechanism the only method whereby ethane, which was observed, might be produced by decomposition of ethylamine is by reduction of the amine. That this reaction
FIG.I does not occur may be inferred from the absence of any effect of added hydrogen found by Taylor. Evidence also is presented in the following work that the effect of added hydrogen on the decomposition of propylamine is no more than that of a similar amount of helium or nitrogen. It seems preferable to assume then in the case here studied that propane and ammonia would be formed by reactions subsequent to the alternative split into propylene and ammonia, particularly since the temperatures here used are considerably lower than those employed by Hurd, from whose data even the percentage of unsaturated hydrocarbon is largely increased by a decreased time of contact. The apparatus for and method of determination of the rate of decomposition were similar to those used previously by Taylor, with the exception that the reaction vessel was of quartz sealed by a graded seal to the pyrex manometer capillary tubing. The propylamine was redistilled from a n Eastman sample and boiled a t 48.j-49.j°C. During the reactions only small amounts of a black tarry deposit were observed in the capillary tubing just above the furnace. The reaction bulb itself was always perfectly clean. The decomposition was studied a t four temperatures a t 5 2 0 , 540, 560 and 580'C. The results are shown graphically in the accompanying figures, the
H. AUSTIS TAYLOR Ah'D HAROLD E. ACHILLES
2660
pressure increase being plotted against time for different initial pressures. The pressure increase was on the average only slightly greater than the initial pressure, the ratio of the total pressure to the initial pressure lying between 2.1 and 2.2. This additional pressure increase is easily accountable on the basis of subsequent reactions. Actual tests made with propylene itself always showed decomposition evidenced by an increase in pressure. To demonstrate the homogeneity of the reaction, quartz tubing was added to the reaction bulb sufficient approximately to double the surface.
FIG.2
Table I contains comparative results between determinations carried out in the empty bulb and with increased surface. Data obtained using a pyrex bulb are included.
TABLE I Temperature j 8 c T Initial Pressure Time in mins.
Quartz Bulb
Double Surface
Pyrex Bulb
209
206
213
AP
AP
AP
0.j
80
a1
80
1.0
40
39
40
1.5 2.0
24
22
23
14
3.0
20
I8 I9
4.0
10
16 =7 9
6.0 8.0
9
IO
IO
I1
6 8 7 The similarity of the rates of pressure change is indicative of the absence of any heterogeneity.
THE THERMAL DECOMPOBITION OF PROPIZASIINE
FIG.3
2661
2662
25,
H. AUSTIN TAYLOR AND HAROLD E. ACHILLES
The order of the reaction was inferred from a n inspection of the times of so and 75 per cent decomposition. Table I1 gives these values.
TABLE I11
Initial Pressure
t2S
Temperature
208
4.2 3.7 3.6 3.35 3.6 3.6 2.82
211
2.70
307 3'0 362
2.46 2.44 2.49
36 38 76 +.*
I /
I03 I Oj
t30
t7j
mins.
s20°C. 10.2
18.9
9.2 7.8 7.3 8.8 8.6 6.5 6.1 5.5 5.4 6.3
20.0
16.4 15.2
17.4
17.5 16.0 14.7 12.7 1 2 .o
-
Temperature 54oOC. 10.5
2.0 1.7
4.9 4.0
8.8
1.j
4.0
10.0
.36 1.10
3.1 2.65 2.82
I
1.10
7.3 6.2 7,4
Temperature j6ooC. 43 79 104 201
301 3 46
0.90 0.90 0.87 0.68 0.58
2.31 2.30 1.79 1.53
0.j 2
I . j;
2.27
5.35 5.30 5.07 4.29 3.60 4. I O
Temperature 58ooC. 39 80 102
203 305 373 432 j02
j68 597
0.j o 0.48 0.43 0.35 0.31 0.31 0.32 0.32 0.31 0.32
I . 16 1.33 I .oo 0.90 0.80
0.81 0.85 0.81 0.80 0.83
2.70
3.30
2.60 2.44
1.95 -
-
The approximate constancy of these values over large pressure changes can only be significant of a reaction which is practically wholly unimolecular. The striking feature is that the values appear to decrease with an increase of
THE THERMAL DECOYPOSITION O F PROPYLAMINE
2663
initial pressure, a decrease which might be considered steady if measurements had not been made a t higher pressures. A close analysis of the data at j80°C., a t which temperature measurements were made a t higher pressures than were used a t the other temperatures, reveals the fact that the values of t 2 5 and tm are constant over the pressure range from 2 0 0 to 600 mms. but that below 2 0 0 mms. the times increase with a pressure decrease. Such a behavior is similar to that found in other unimolecular reactions although it should be noted that the rate of falling off a t lower pressures is very considerable slower than is observed in the decomposition of the ethers. The times of three-quarter decomposition do not exhibit the same degree of constancy that the other values show, presumably since after that time the amounts of secondary decompositions such as of propylene or ethyl cyanide are appreciable in comparison with the primary reaction, the order of these subsequent reactions being uncertain. That this explanation is substantially correct is seen from the effects of foreign gases on the reaction as a whole, where for example a t smaller pressures of added diluent an effect only becomes noticeable towards the end of the reaction. In Table I11 are presented the rates of pressure change observed using various amounts of hydrogen and nitrogen added initially to the amine.
h B C .4
B
= = =
T a B L E 111 initial pressures of amine initial pressures of hydrogen initial pressures of nitrogen 210 302
406
202
51 73
45 65
Time 0 .j
I
.o
1.5
85
2.0
93
76 82
3 .o
IO0
88
4.0
105
6.0
I08
93 96
8.0
I10
6
I99 207*
0.5 .o I .j 2 .o I
3.0 4.0
6.0 8.0
* Helium.
56 88
98 202 302
50
200
400
41 65
'09
76 89
I22
IO1
85
137 146
109
93 97
I I j
76
15.2
I20
I02
156
123
105
2664
H. AUSTIN TAYLOR AND HAROLD E. ACHILLES
It can be seen therefore that at the lower pressures of foreign gas no effect is observed until the decomposition is over 7 5 per cent complete. Further, the effects of hydrogen and nitrogen are similar, except when large excesses are used, when hydrogen has a somewhat greater effect than nitrogen. Finally the effect of helium is seen to parallel that of nitrogen at the same pressure. The absence of an effect of added diluent at lower pressures corroborates the previous evidence for the unimolecular nature of the reaction. The effects in slowing the rate of reaction at higher pressures of diluent might be indicative of inelastic collisions with activated propylamine molecules. Such a behavior by hydrogen however is contrary to the general results obtained in numerous other unimolecular rractions where hydrogen is generally found to have the effect of increasing the rate. The alternative and more probable explanation of the effect of the diluents in decreasing the rate must lie in the secondary decompositions subsequent t o the primary unimolecular split of the amine. These secondaryreactions will be at least bimolecular, the probability of their beingunimolecular at least early i n the reaction is small since they are then present only in small amounts, The effect of added gas therefore would be to reduce the rate of these reactions. That perceptible amounts of these react,ions do occur even early in the reaction can be seen from the falling value that is obtained for t,he velocity constttnt a n example of which is given in Table IT. TABLE IV Teniperature Time
Pressure Change
k
58oOC'.
Time
1I 9
0.729
I .o
18j
0
653
5.0 8.0
1.5
231
0 600
2 . 0
259
3.0 4.0
0 . j
Pressur(< Change
k
340.5 3.;'. j
0,344 0.308
I2.C
365
0.229
0.546
16.0
0.196
296
".jj5
24.0
371 378
318
c.423
0,139
The constants were calculated using the extrapolated limiting pressure increase as proportional t o the initial concentration of amine. If the secondary reactions are inhibited by added gases the values calculated for the velocity constants from the pressure change in presence of added gases should be more nearly constant. This is actually the case. Although the constants still show a decided drift, the extent of it is by no means as great as that above. Regarding the energy of activation, since, it will be recalled, the quarter and half lives are not constant over the full pressure range studied, the logarithms of these values for a given initial pressure were each plotted against the reciprocals of the absolute temperatures. In this way graphs were obtained at initial pressures of about 40, 80, 100, 2 0 0 and 300 mms. The graphs were in all cases straight lines and the parallelism of the lines may be judged from the similarity of the values of the energies of activation calculated from their slopes and given in Table V.
2665
THE THERMAL DECOMPOSITION OF PROPYLAMINE
TABLE V
Initial Pressure
Initial Pressure
calories
40
44J6Oo
2 00
44,900
80
4315"
300
44,900 44,400
IO0
E in
43,900
E in calories
Mean
This value is quite similar to the value found by Taylor for the energy of activation of ethylamine namely 43,400 calories and may be significant of a similar primary rupture of the molecule in the two cases. One important feature of the above results is the identity of the energies of activation when calculated for the same initial pressure a t different temperatures, proving that the rate a t which the reaction deviates from the truly unimolecular course is the same a t different temperatures and further that the pressure below which the reaction is no longer unimolecular is approximately the same over this temperature range. A comparison of the above data with those of other unimolecular reactions may be made by means of the ratio of E/RT a t temperatures where the reactions have equal velocities. The rather low value of 2 8 . 2 ' was found by Taylor a t joo°C. for the ethylamine decomposition, whilst from the above quarter and half lives propylamine can be shown to be decomposing a t the same rate a t approximately the same temperature and would therefore give a value of 28.9. How relatively low is the energy of activation of propylamine can perhaps best be seen from a comparison of the ratio of effective to total collisions occurring and the value of e-E RT. Taking a velocity constant of 0.00158 sec.-l a t 5oo°C. the number of molecules reacting per cubic centimeter per second is approximately 4 X 1 0 ~ ~ The . reaction deviates from its unimolecular course a t about zoo mms. a t which pressure therefore, (using a molecular diameter of 7 X IO-^ cm.) there will be 3.5 X 1 0 2 7 collisions per cubic centimeter per second. The ratio is therefore I X 1 0 - l ~approximately. Using the value 44,400 calories for the energy of activation derived from the simple Arrhenius equation the value of e-E'RT becomes about 3 X 10-l~. Keglecting then, the several small corrections which might be made in these figures there is seen to be a substantial agreement. By analogy with other unimolecular reactions this agreement would suggest that a single vibrational bond, corresponding that is, to two square terms, is alone responsible in the activation process, and since the data for ethylamine and propylamine are so similar it may be concluded that the same bond is involved in both cases. X comparison of the data for the decompositions of dimethyl, diethyl and dipropyl ethers, as also for azomethane and azoisopropane shows marked differences with increasing complexity of the molecules quite unlike the similarity found here between ethyl and propyl amines. X possible explanation suggests itself in that in the ethers and azo compounds the primary rupture probably This value was given in error in the original as 28.9.
2666
H. AUSTIN TAYLOR A S D HAROLD E. ACHILLES
occurs deep within the molecule whereas with the amines the rupture is most probably in the C-X bond or at least at the end of the carbon chain. Further work is in progress with more complex amines to determine more precisely this influence of complexity. Sight should not be lost of the fact, however, that the secondary reactions for the amines are extremely complex and that consequently the energies of activation may also be complex.
Summary The decomposition of propylamine over the temperature range j2o-58o0C. is shown to be a homogeneous reaction. At pressures above 2 0 0 mms. the primary reaction is apparently unimolecular. An energy of activation of 44,400 calories shows similarity to that found for ethylamine. ,Yichols Chemical Laboratory, York Universzty, S e w Y o r k , X. Y .
Seu;
