THE THERMAL DECOMPOSITION OF DIMETHYLAMINE BY H. AUSTIN TAYLOR
Previous kinetic studies of the decompositions of ethylamine,' of propylamine2 and of i~opropylamine~ have shown that although the reactions are complex it is reasonably certain that the initial process is homogeneous and unimolecular in each case. The observed decrease in the rates of decomposition for propylamine and isopropylamine, with decreasing pressure led to the conclusion that activation of a single vibrational bond alone was involved. Such an apparent simplicity is to be contrasted with the very complex mechanisms involved in the activation process of other unimdecular reactions. In an effort to see what effects a secondary amine would have on such results the present investigation was carried out and although appearing on the surface to be very similar, its extreme complexities suggest that the simplicity of mechanism of activation is probably fortuitous due to a mutual compensation of several reactions. Further work on the amines is still in progress but it seems advisable that a report of the type of complications involved should be given since a similar condition may possibly account for certain peculiarities in other cases. The apparatus and procedure were identical with those already published. The results at pressures below 15 mms. were obtained by the use of a sloping manometer whereby the pressure change was multiplied by a factor of 8.26. The dimethylamine was redistilled from an Eastman sample and showed a small trace of a primary and tertiary amine. The decomposition was studied over the temperature range 480-51o'"C. at pressures from 3 . 7 5 to 600 mms. The pressure increase during reaction averaged I 20 percent although a slight drift with pressure was noticeable, the increase being greater a t lower pressures. No effect of temperature on this value could be observed. The increases in pressure plotted against time give typical unimolecular rate curves showing, however, a very slow but continued pressure increase towards the end of the reaction. I t was necessary therefore to follow a reaction which for example had a quarter life of 3 minutes, about 15 hours to obtain a reliable end-point. I n Table I are given the quarter lives for the various temperatures and pressures studied, calculated as the time taken for the pressure to increase by one-quarter of the total change. The natural conclusion from such a constancy of the quarter lives at the higher pressures is that the reaction is unimolecular. At lower pressures the order of reaction would appear to be between one and two.
a
J. Phys. Chem., 34, 2761 (1930). J. Phys. Chem., 35, 2658 (1931). J. Phys. Chem., 36,670 (1932).
1961
THERMAL DECOMPOSITION OF DIMETHYLAMINE
TABLE I Initial Pressure in mms.
t a s mins.
Initial Pressure in mms.
Temperature 480°C. 9.9 279 8.8 333 7.8 597
88 I33 181
I20
Temperature 49Ooc. 6.5 229 5.6 303
176
5.2
81
5.8
4.6 3.8 3 .o
555
2.8 3 .o 2.4
10.0
53 105
116
mins.
5.3 4.6
Temperature 5ooOC. 161 8 .o 198 5.4 251 3.5 315 3.1 547
5.8 IS
tab
2 2 2.1
Temperature ~IoOC. 3.75 14.8 44 90 I 60 I73
7.8 4.9 3.6 2.6
225
2 I
297 404 409 42 1 486 522 524 622
2.2
2.2
252
.o .8
Since a change in end-point had been observed with change in initial pressure it seemed possible that the reaction might be in part a t least heterogeneous. On increasing the surface of the pyrex reaction vessel by the addition of measured lengths of pyrex tubing whereby the surface to volume ratio was increased 3.4 times, the quarter lines given in Table I1 were obtained.
TABLE I1 Increased Surface (S/V = 3.4) Temperature Initial Pressure
4 41 161
~ I O O C .
tas
8.0
4.0 2.2
These values, it will be seen, are in substantial agreement with those previously quoted showing the absence of any marked heterogeneity.
1962
H. AUSTIN TAYLOR
Assuming then the react'ion to be homogeneous and unimolecular it will be seen that at groOC. the velocity constant appears to fall off below about 300 mms. I t is of interest to see the effects of various added gases. The effect of the reaction product's was first' investigated. The decomposition of 1 2 0 mms. of dimethylamine was allowed to go to completion. A furt,her charge of amine was then added and t>herate followed in the usual manner. Two such runs were made in t>hepressure region where the reaction rate had fallen off namely at 94 and 44 mms. each with the products from 120 mms. of the amine present,. The quarter life was 2.0 minutes in each case as compared with 2.6and 3.6 minutes, respectively, obtained in the absence of products. It appears, therefore that the products are capable of maintaining the rate somewhat in the lower pressure range. A similar effect was observed with hydrogen, ethane and ammonia. With zoo mms. of hydrogen the quarter lives of 39 and 161 mms. of amine were 2.4 and 1.6 minutes, respectively, compared with 3.9 and 2 . 2 minutes without hydrogen. A similar amount of ethane gave times of 2.8 and 1.6 minutes for 47 and 162mms. of amine whilst zoo mms, of ammonia gave for 46 and 164 mms. of amine quarter lives of 2.8 and 1.6 minutes. In all the reactions involving either t,he products, hydrogen or ethane the end-point of the decomposit,ion was the same as in the absence of added gases. With ammonia, however, the total pressure increase rose from 1 2 5 to 135 percent of the initial amine pressure which might suggest some actual chemical interaction although even t,his extent of change is not, as great as the I jo percent observed with 4 mms. init,ial pressure of amine alone. With et'hylene as added gas more definite evidence of reaction was obtained. Even making due allowance for the pressure decrease due to polymerisat'ion of the ethylene, dat'a for which were obtained from a blank run with ethylene alone, the observed pressure increase due to the amine was only 80 percent of the initial pressure. The general course of the reaction too was quite changed notably with respect to the end-point which was reached very sharply after only 40 minutes in comparison with the several hours necessary with amine alone. A comparison of the rate of reaction in presence of ethylene with that in its absence is thus impossible. One definitely inert gas was tried namely nitrogen and found to be without effect at any stage of the reaction. When zoo mms. of nitrogen were added to 161 and 485 mms. of amine quarter lives of 2 . 2 and I . j minutes were obtained in complete agreement wit,h the values obtained for the amine alone. Such a behavior by a reaction would appear to be best explained on the assumption that the reaction is unimolecular. The constancy of rate a t higher pressures, its falling value at lower pressures, the absence of effect of truly inert gases at any pressure, the maintenance of the high pressure rate in the low pressure region by certain complex molecules with absence of effect by the same molecules in the high pressure range are all usually accepted as unequivocal evidence of unimolecularity in a reaction. The react,ion rates obtained may thus be compared with others similarly obtained. The energy of activation obtained by plotting the logarithms of the limiting quarter lives at higher pressures against the reciprocals of the
THERMAL DECOMPOSITION OF DIMETHYLAMINE
1963
corresponding absolute temperatures and taking the slope of the straight line so given is found to be 44,300 calories. The value is identical within the limits of error with that obtained for ethylamine and for propylamine. With this value one can calculate that a reaction would occur with a velocity constant of 0.00158 sec.-l corresponding to a quarter life of 3.0 minutes at 489% At this temperature the value of E / R T would be 29.3 as compared with 28.2 obtained for ethylamine and 28.9 for propylamine. Taking the falling off as occurring around 400 mms. a t this temperature, the number of molecules entering into collision under these conditions is 2 . 7 X IO**per cubic centimeter per second using a molecular diameter of 7 X IO-^ cms. The number of molecules reacting similarly is 4 X 1015 giving a ratio of 1.5 X 10-l~. The value of is 1.8X1o-l~an agreement which would only be possible with a molecular model involving one or two effective degrees of freedom, a condition similarly found for the other amines studied. Despite the apparent consistency between all the amines so far studied there is certain evidence that the results may be misleading, certainly in this case with dimethylamine and possibly in the previous ones also. It seems hardly conceivable that an extremely complex reaction could yield reaction rates which on the surface appear to agree so well with accepted theories of unimolecular reactions. The percentage pressure increase during the reaction could well be accounted for on the basis of a single molecule yielding two molecules on initial decomposition followed by a small further decomposition of the products. Such a split in the case of dimethylamine might be presumed to yield ammonia and ethylene. An analysis of the products of decomposition after one of the above runs showed no ammonia present and no unsaturated hydrocarbon. The quantity of material available for analysis in any single experiment as actually carried out in a static manner being small, a definite weight of the dimethylamine was sealed in an evacuated glass bulb and placed in the furnace for a sufficient time for reaction to be complete. The bulb was then cooled and the contents analysed. The appearance of the bulb before opening was surprising, for despite the fact that the vessel used in the static experiments was perfectly clean after several hundred runs and that only a small amount of a dark brown liquid had collected in the capillaries just outside the furnace, the bulb was black and quite opaque to light. It should be noticed, however, that the actual concentration in the latter case was very considerably greater than in any one of the static experiments. The bulb was opened over water in such a way that the insoluble gases could be collected and measured. For about an hour after the bulb had been opened a contraction in the volume of the system was observed pointing to a slow solution of some constituent. Analysis of the fixed gas showed it to contain 65 percent methane, 30 percent ethane and 5 percent hydrogen, no free nitrogen being present. An acid titration of the aqueous solution yielded only 0.4 of the nitrogen initially present in the amine taken. The bulb still carrying its black coating was next examined. The odor was entirely different from that of an amine, being almost putrid. The addition of hydrochloric acid, however, caused most of the black coating to dissolve as a dark brown solution carrying
I964
H. AUSTIN TAYLOR
a small amount of a black solid in suspension, presumably carbon. The amount of this latter acid addition made up almost completely for the nitrogen content of the amine taken. By subtracting the amount of hydrogen and carbon found from the initial amine taken an empirical formula for the residue of CHsNs was found. Though probably not guanidin, as t,his might suggest, the final product must be some complex nitrogenous material of that type. In an effort to isolate any intermediate substances that might be formed during the react,ion several dynamic runs were made. The dimethylamine was allowed to evaporate under its own vapor pressure a t room temperature through a small capillary into a furnace kept at 50o'C. The products obtained consisted of a colorless liquid with strong ammoniacal odor but penetrating after-effect suggestive of a hydrazine. I t boiled over a range from about 7 0 to IIoOC., reduced silver oxide in the cold and decomposed at ;oo"C. in a static run wit,h a pressure increase at a rate almost ten times as great as that of the dimethylamine. The liquid contained a small amount of a white solid, soluble in water but insoluble in ether, which sublimed without melting at about z jo"C. and gave a dull orange precipitate with iodine in potassium iodide. Analysis of the fixed gases showed 60 percent of hydrogen and 40 percent, of ethane. The aqueous solution over which the gases were collect,ed showed unchanged secondary amine and a trace of ammonia shown by the insoluble sodium ammonium cobaltinitrit,e. I t seems probable, therefore, that the products contain chiefly methyl hydrazines, the solid being probably hexamethylene tctramine. On raising the temperature of the furnace to 6oo°C.,using t'he same rate of flow of amine as previously, the liquid which collected mas no longer colorless but. decidedly brown, its complexity being shown by its boiling point range there being some constituent present decidedly more from about, 40 t,o I~OOC., volatile than in the liquid previously obtained at the lower temperature. This more volatile constituent decomposed at j0o"C. in a stat'ic experiment with a pressure decrease at a rate which was about five times as rapid as the rate of pressure increase of the original amine. Furthermore, the brown liquid maintained at 100°C. gradually developed larger quantities of the more volatile constituent decomposing wit,h a pressure decrease. This pressure decrease was undoubtedly due to a polymerisat,ion to a high boiling product which subsequently condensed in the capillaries outside the furnace, as considerable difficulty was experienced in evacuating the system after these experiments. The decomposition of dimethylamine under the conditions of the static experiments t,herefore must be extremely complex and the observed rate of pressure increase, a composite of at least t,wo rates in opposing directions. I t would appear that methyl hydrazines are the chief constituent, in the lower temperature dynamic runs and presumably therefore are easily formed from dimethylamine. Their formation should occur as a bimolecular reaction. If the subsequent decomposition of the hydrazine was unimolecular one would expect an induction period in the net rate unless the primary bimolecular reaction was much more rapid than the unimolecular. I n this latter event a
THERMAL DECOMPOSITION OF DIMETHYLAMINE
I965
lowering of temperature might possibly permit a separation of the two. No such change could actually be found. This of course may simply mean that the energies of activation of the bimolecular and unimolecular rates are approximately the same.1 In any event the very rapid rate of decomposition of the hydrazines is being offset by the reaction with the pressure decrease. It seems almost phenomenal that such a complex reaction should yield what are apparently simple data and one naturally wonders whether it is possible to reproduce for example pressure-rate curves such as are being accepted as typical of true unimolecular reactions by a counterbalancing of effects in two or more reactions proceeding simultaneously. It may be shown in certain cases that a bimolecular polymerization proceeding simultaneously with a unimolecular decomposition may give a logarithmic relation between initial pressure and half-life as calculated from the observed pressure change. Should the form of this logarithmic relation be such as to bring the asymptotic portion of the curve within easily measurable pressure ranges an apparently constant half-life with changing initial pressure would be observed and might be mistaken for a true unimolecular reaction whose rate decreased a t lower initial concentrations. Whether such is the explanation of the case here studied remains for further work to prove. One rather significant feature of the resulfs obtained, lies in the rapidly falling values calculated for the velocity constants in any run, despite the fact as shown that the products of the reaction tend to maintain the high pressure
TABLE I11 Temperature 5 I o O C . Time
AP
I
125
2
205
3 4
269 314 351 407 44 1 469 489 503 527 543 5 53 567 578 587
5 7 9 I1
13 I5 20 25
30 40 50
60
Initial pressure 486 mms. ki ka X 10s .217 .I93 .182 .169 'I59 ' I44 ,130 *
I20
.I11
.103 .087 .076 .067 ,054 .047 .042
.379 .368 .378 .376 ,380 ,390 .390 ,390 .389 .382 .364 .350 .33I .303 .zgr .z88
l A caae in point is actually found in the decomposition of ethyl mercaptan to be published shortly.
1966
H. AUSTIS TAYLOR
rate. Table 111 gives a typical example of the calculated velocity constants, kl being the unimolecular constants, kz the bimolecular constants. I t will be noted that the values of kz are much better than kt although from the quarter lives the reaction cannot possibly be bimolecular. In conclusion one further point may be raised. The accompanying diagram shows the curves obtained by plotting the quarter lives against initial
I
I 70
1%
a 0
280 so Peeasme I
420
'190
5th
630
I"".
FIG.I pressures at the four temperatures studied. The forms of the curves might well be those of a unimolecular reaction were it not for the fact that the reaction rate appears to decrease a t a higher pressure the lower the temperature. Such a variation is contrary to the usual expectation for a unimolecular reaction. A plot of the quarter life against the reciprocal of the initial pressure gives at each temperature curves showing such a marked curvature that an extrapolation to determine the limiting quarter life at high pressures is doubtful but would indicate that the limiting value has not been reached at 600 mms. at the lowest temperature studied. Summary The rate of pressure increase in the deconiposition of dimethylamine over a temperature range from 480 to 510°C. appears to indicate a homogeneous unimolecular reaction with an energy of activation of 44,300 calories. Dynamic investigations show the reaction to be so extremely complex as t o cast doubt on the above interpretation. Further work is in progress. Wzchols Chemzcal Laboratory,
New York Unwerszty, New York, N. Y.
