T H E THERMAL DECOMPOSITION OF DIETHYLAMINE' H. AUSTIN TAYLOR
AND
C. R. HERMAN
Nichols Chemistry Laboratory, N e w York University, N e w York Citg Received November 84, 19.34
The study of the kinetics and mechanism of the decomposition of diethylamine was undertaken in order to continue the work in the amine series under investigation in this laboratory. The decompositions of the primary amines appeared a t first sight to offer simple examples of homogeneous unimolecular reactions (8,9,11). This simplicity, however, was shown by Taylor (10) to be illusory in the decomposition of dimethylamine, owing to a mutual compensation of several concurrent reactions. It was hoped therefore that a fuller investigation, particularly of the mechanism, might yield a possible scheme of analysis of the complete results. In a study of the catalytic influence of iodine on amines, Bairstow and Hinshelwood (1) assume the over-all change occurring with diethylamine to be (C2Hs)zNH --j CzHe
+ CHd + HCN
despite the fact that in the uncatalyzed reaction only from 15 to 30 per cent of the theoretical amount of hydrogen cyanide could be detected with silver nitrate and that this fell to from 5 to 10 per cent in the iodinecatalyzed decomposition. They conclude that the reaction is initially homogeneous and unimolecular. In view of the fact, as will be shown, that methane and nitrogen are the chief products and that no hydrogen cyanide was detected in any stage of the reaction, although a hydrazine which readily reacted with silver nitrate was always found, this mechaiiism appears untenable. Furthermore, there appears a distinct possibility that hydrogen cyanide, if formed, would add on to any unchanged amine and thus reduce still further the observed rate of pressure increase. The rate of decomposition was studied statically in an apparatus identical with that used in the previous work already published. The diethylamine used was an Eastman sample, redistilled over lime, and boiling from 55.5 to 56.0"C. The capillaries connecting the amine reservoir, reaction vessel, and manometer were maintained a t 60°C. to prevent any condensation occurring. Data were obtained over the temperature range Abstract from a thesis presented by C. R. Herman in partial fulfillment of the requirements for the degree of Doctor of Philosophy a t New York University. 803
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H, AUSTIN TAYLOR AND C. R. HERMAN
from 510 to 540°C., and a t pressures from 40 to 400 mm. The pressure increase during reaction averaged around 160 per cent, although a perceptible effect of both temperature and pressure was observable. Thus at 510°C. the pressure increase varied from 182 to 166 per cent at pressures ranging from 44 to 419 mm., whilst a t 540OC. the variation was from 176 to 134 per cent over the same pressure range. Such variations, however, may not be a true indication of the existence of an equilibrium but may be
5
10 TIME IN MINUTES
FIG.1. THERMAL DECOMPOSITION OF DIETHYLAMINE
merely due to the fact that the end points taken are fictitious, owing to very slowly occurring secondary reactions. A typical diagram of the rate of pressure change with time is shown in figure 1 for initial pressures of 44, 91.5, 156.5, 203, 291.5, and 401 mm. at 510°C. Similar curves were obtained at the other temperatures. In table 1 are given the values of the quarter-lives for the various temperatures and pressures studied, calculated as the time necessary for one quarter of the total pressure increase to occur. The only obvious conclusion one can draw from these vaIues is that the reaction has an order lying between one and two, more closely approaching
805
THERMAL DECOMPOSITION OF DIETHYLAMINE
the former however, but being in all probability complex. The variation in quarter-life with pressure is the more marked the lower the pressure, and the possibility therefore remains that a t higher pressures a constant life might be found. The reaction was shown to be homogeneous under the conditions studied by increasing the surface to volume ratio of the reaction vessel eleven times by packing with short ,lengths of Pyrex tubing. Table 2 shows a compariTABLE 1 Values o j the quarter-lives at different temperatures and pressures INITIAL PRESSURE
I
tINYINUTES
Temperature = 510°C. 44.0 58.5 81.0 91.5 156.5 203 257.5 291.5 322.5 374.5 401.0 419.0
3.80 3.35 3.05 2.80 2.23 2.08 1.96 1.87 1.74 1.66 1.56 1.50
Temperature = 520°C. 37 85 127 142 185 204 245 282 385 432
'
2.30 1.87 1.65 1.48 1.35 1.30 1.21 1.15 1.03 0.96
INITIAL PRESSURE
I
1INYINUTES
Temperature = 530°C. 43 71 88 134 188 228 256 295 342 368 410 452
1.45 1.34 1.15 1.09 0.92 0.88 0.83 0.76 0.74 0.74 0.70 0.69
Temperature = 540°C. 41 91 129 152 209 252 275 315 358 439
1.20 0.85 0.77 0.68 0.63 0.60 0.57 0.56 0.54 0.50
son of the actual pressure readings at various times for the same initial pressure in the packed and unpacked vessels. The data in table 2 show that the reaction is in the main homogeneous, though there are indications of a speeding up over the later portion of the reaction, though the pressure increase for the unpacked vessel was 174 per cent whilst that for the packed vessel was 175 per cent. The effects of added nitrogen, hydrogen, and ammonia were studied at the highest and lowest temperatures, namely, 540 and 510°C., and a t
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H. AUSTIN TAYLOR AND C. R. HERMAN
initial pressures of amine of 40 and 200 mm., the pressure of added gas being 150 mm. These data are presented in table 3 in the form of percentage total pressure increase and the quarter-life. It is readily seen that the effect of nitrogen is negligibly small at all temperatures and pressures. Ammonia, however, appears to retard the rate of pressure increase, and in all probability does so by specific chemical reaction. The effect of hydrogen is very marked. I n the first place the end point is considerably reduced, showing again a probable chemical action. The quarter-life based on this end point is also largely reduced though, owing to the disproportionate end points, the effect is apparently larger than it actually is for the earlier portions of reaction. The actual TABLE 2 Changes in pressure observed when the reaction occurs i n packed and in unpacked veasels Temperature = 510°C. PRESSURE CHANGE
TIME
Unpacked vessel. Initial pressure 233 mm.
Packed vessel. Initial pressure 230 mm.
26 51 99 139 170 200 223 242 258 283 302 316 326
25 50 94 134 168 195 220 240 255 283 303 319 332
minutes
0.5 1 .o 2.0 3.0 4.0 5.0 6.0 7.0 8.0 10 .o 12.0 14.0 16 .O
pressure change for equivalent times is not however as great for the amine in presence of hydrogen as in its absence. There appears then to be an actual retardation by hydrogen. With the view of determining the end products of reaction a weighed amount of amine was sealed in a glass bomb and heated for three days a t 500°C. to render reaction complete. The gaseous products were then analyzed,%yielding 2.1 per cent of ammonia as water-soluble, 3.2 per cent of unsaturated hydrocarbon absorbed by bromine, 0.2 per cent of hydrogen by preferential combustion over copper oxide, 70.1 per cent of methane,
* Thanks are due t o H. Tarnpoll of this laboratory for the gas analyses.
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THERMAL DECOMPOSITION OF DIETHYLAMINE
and 2.3 per cent of ethane, The remaining 23 per cent is nitrogen. The reaction in the main then appears to yield methane and nitrogen in the ratio of three to one. The difference between this value and their ratio in the original amine is to be accounted for by the large amount of black deposit obtained in such bomb experiments. Such results, though of little use in determining a mechanism, are necessary to the elucidation of the final goal of the reaction. To obtain the details of the mechanism of the reaction in its earliest stages a small amount of liquid diethylamine was sealed in a tube and
PRESSURE OF AMINE
I
TABLE 3 The effect of added gases GAS
I
PRESSURE OF GAS
I
PER CENT INCREASE
I
t
Temperature = 510°C. mm.
44 40 43 46 203 201 210 180
minutes
mm.
153.5 151 152 153 152 155
182 179 182 170 175 171 174 157
3 .SO 4.20 3.75 3.55 2.08 2.20 2.17 2.10
176 177 175 166 156 164 156 139
1.20 .1.30 1.17
Temperature = 540°C.
41 41 48 38 210 207 21 2 225
154 148 151 154 153 158
1.10
0.63 0.66 0.65 0.57
maintained at only 200°C. for twenty-four hours. Upon opening, a small pressure increase was noted, pointing to the accumulation of gaseous products. In other cases a t 400°C. and for periods varying from one to eighteen hours similar gaseous products were found and analyzed. In no case was hydrogen found in other than mere traces. Nitrogen was always present. From the ratio of the carbon,dioxide to the water formed on combustion the gas appeared to be a mixture of methane and butane. To test this specifically, a sample of the gas was kept in a solid carbon dioxide bath for some time, during which a decrease in volume was observed. The remaining uncondensed gas on combustion proved to be practically
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H. AUSTIN TAYLOR AND C. R . HERMAN
pure methane. The condensed portion upon vaporizing and combustion analyzed as more nearly pure butane. Analysis of the remaining liquid in these experiments showed a considerable amount of unchanged secondary amine with only traces of primary and tertiary amine. Specific tests for the cyanide group as, for example, by the Prussian blue test showed its complete absence. In like manner nitriles too were shown to be absent. Upon addition of silver nitrate a marked reduction occurred, yielding the characteristie silver mirror. It is probable that this reduction is caused by a hydrazine, and the initial reaction may tentatively be indicated as
2(CzHs)ZNH + C2HsNH-NHC2Hs
+ C4H10
a bimolecular reaction involving no volume change. Subsequent reaction involving the liberation of nitrogen from the hydrazine and the decomposition of butane would account for the presence of nitrogen, methane, and ethane in the end products, the over-all volume increase, and for the observed slow secondary reactions yielding apparent equilibria in the final state. The absence of hydrogen in the high temperature bomb experiments is in line with the observations of rehydrogenation of unsaturated hydrocarbons in the butane decomposition (2,3). The excess of methane over ethane is to be expected also because of decomposition of the ethane (4). The decomposition, at least in its earlier stages, of butane, is admittedly a unimolecular reaction even if of a chain type ( 5 ) ) and it would appear reasonable from the complexity of the molecule that the diethylhydrazine would also decompose unimolecularly . This is under investigation in this laboratory a t the present time. Assuming that this latter rate is faster than the observed rate of pressure increase of diethylamine itself, the rate actually being measured here would be that of the bimolecular sp1,it into butane and hydrazine, a reaction without a volume change, in terms of the volume increase of the unimolecular reactions. It does not seem advisable a t present to speculate on these relative rates, but the observation that the rate of pressure increase found would indicate a reaction order between one and two is easily accounted for, other than by the assumption that it is due to a quasi-unimoleoular reaction in the pressure range where the Maxwell-Boltzmann distribution of activated molecules is not maintained. I n view of the fact that all analyses of the reaction products, whether taken in the early stages or later on, show the presence of some primary amine, ammonia, and unsaturated hydrocarbon, the possibility is suggested of an initial split of the amine into ammonia and an unsaturated hydrocarbon. The quantities of the latter usually found are hardly more than traces, however. Such a rupture therefore must be much more difficult than the hydrazine formation and consequently occur only to a minor ex-
\
THERMAL DECOMPOSITION OF DIETHYLAMINE
809
tent. The reverse reaction, between ammonia and the unsaturated hydrocarbon, will account for the repressing effect of the added ammonia previously mentioned, since in the butane decomposition some unsaturateds must certainly be formed, as may also be possible in the hydrazine decomposition. The effect of added hydrogen would be similar though more marked. There remains one factor, following as a consequence of the suggested mechanism, to be discussed. If the major reaction is bimolecular and without a volume change, followed by unimolecular reactions, there should be expected a period of induction in the reaction. Depending on the relative rates this may be long or short. However, as was first pointed out by Schumacher and Wiig (7) in the case of ethylamine, the induction period only evidences itself on an extremely clean surface and a t low initial pressures, and even then appears to show only very poor reproducibility. Under the present conditions the reproducibility is perfect even with an extended surface, though the argument might. be advanced that the surfaces were all uniformly poisoned. Under such conditions it is legitimate nevertheless, to treat the reaction as homogeneous. There seems however no reason to doubt that if the decomposition of diethylamine were to be studied in a vessel which had been pumped out to less than mm. for several days, an induction period would be observed, as is the frequent observation with ethylamine in this laboratory a t present. Indications in this case seem to point however to the hydrazine formation reaction as the one which is so extremely sensitive to traces of foreign material. An approximate idea of the energy of activation of the early reaction may be obtained from the observed quarter-lives. The average value yielded by the simple Arrhenius equation for the various pressures studied is 49,000calories. This value, though somewhat higher than that found in a similar manner for the primary amines, namely 44,000 calories, is more nearly in agreement with the value of about 50,000 calories taken as indicative of the strength of the C-N bond by Rice and Johnson (6) from a study of the free radical formation by thermal decomposition. SUMMARY
The rate of pressure increase in the diethylamine decomposition over a temperature range from 510 to 540OC. appears to indicate a homogeneous reaction with an energy of activation, in the early stages, of 49,000calories. Analysis of intermediate products suggests the mechanism to involve the formation and subsequent decompositionof diethylhydrazine and butane. REFERENCES (1) BAIRSTOW AND HINSHELWOOD: J. Chem. soc. 1933, 1158. (2) FREYAND HEPP:Ind. Eyg. Chem. 26,441 (1933).
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H. AUSTIN TAYLOR AND C. R. HERMAN
(3) PEASEAND DURCAN:J. Am. Chem. SOC. 62, 1262 (1930). (4) PEASEAND DURQAN: J. Am. Chem. SOC.60,2718 (1928). (5) RICEAND HERZFELD: J. Am. Chem. SOC.66,284 (1934). J. Am. Chem. SOC.66,219(1934). (6) RICEAND JOHNSON: AND WIIQ: Z. physik. Chem. 162A,419 (1932). (7) SCHUMACHER (8) TAYLOR: J. Phys. Chem. 34,2761 (1930). (9) TAYLOR: J. Phys. Chem. 36, 670 (1932). (10) TAYLOR: J. Phys. Chem. 36, 1960 (1932). (11) TAYLOR AND ACHILLES: J. Phys. Chem. 36, 2658 (1931).
