T H E DEC03IPOSITION OF ETHYLAMINE. X ITXIMOLECULAR REACTION BY H. ACSTIN TAYLOR
The theoretical difficulties attending the problem of the mechanism of unimolecular reactions demands a t the present time an accumulation of data on such reactions which nil1 permit a definite decision to be made between the various theories that have recently bsen propounded. The studies by Hinshelwood' of the unimolecular disruption of various organic aldehydes, ketones and ethers have furnished considerable data on the breaking of C-C, C-H or C - 0 bonds, whilst the decompositions of azomethane and azo-isopropane by Ramsperger* offer evidence on the C - S bond. I n this latter connection the amines offer a series of compounds similar in structure to the compounds studied by Hinshelwood which should afford useful comparative results. The decomposition of ethylamine here studied is shown to be a homogeneous unimolecular reaction down to pressures of 5 0 mms., occurring a t temperatures similar to those necessary for the deconiposition of the corresponding ketones and ethers, namely 500" to 550' C. The investigation mas carried out statically by observing the rate of pressure change with time. Redistilled ethylamine, boiling a t 18.i oC was maintained liquid by immersion in cold water kept continuously running. The vapor pressure was always sufficient to furnish reactant for experiment. The reaction system was made entirely of pyrex glass, the portion in the furnace being about 300 ccs. in volume. The inlet tube and connecting manometer were of capillary tubing, the total volume outside the furnace constituting about one percent of the total volume of the system. h 'hyvac' oil pump was used to evacuate the system: a platinum resistance thermometer furnishing the necessary temperature measurement. The decomposition of ethylamine involving the production of ammonia and ethylene should result in a pressure increase equivalent to the initial pressure taken. The results indicate this to be true at least at the lower pressures studied, the deviations a t higher pressures will be discussed later. That the reaction is unimolecular in nature may best be judged from the data in Table I which gives the time taken for the initial pressure to increase by 2 j , j o and j percent. There can be no doubt whatsoever that the reaction is essentially unimolecular. The constancy of the quarter life is indisputable in itself whilst the half and three-quarter life is sufficiently constant t o afford corroborative evidence. The actual results obtained are shown graphically in the accompanying diagrams. The forms of the curves are typical of a unimolecular change. Proc. Roy. SOC.,111A, 2 4 j (1926); 113A, 221 (1926); 114A, 84 (1927); 115A, 215 (1927);
J. Chem. SOC., 1929, 1804. * J. Am. Chem. SOC.,49, 912 (1927); Proc. Nat. .4cad. Sci., 13, 847 (1927).
If. A U S T I S TAYLOR
2762
TABLE I Initial Pressure
t25
t ; j mins.
t6,
Temperature
500'
C.
392
3 6
8.4
18.4
318
3.6
8.4
18.0
I i .o
257
3.2
245
3.6
186
3.0
7 .6 8.j .8
'3'
3 6 '
9 0
3 5 3.6
8.8 8. o
55 37
Temperature
17.4 16 . o 19.0 r8.j r.; . o
j ~ o c'. "
212
2 4
6.2
136
2 4
82
2 4
6-1 f> 0
-17 :. ?
5
..j
5.2
j .6
TEE DECOVPOSITION O F ETHI'LAMIXE
2763
To tlemonstratr the.homogencity of the rcnction espcrimcnts were carried wit lritli nn incrcnsed surface in tlir rc'nction vessel. In Table I1 are giver1 coniparntive results hctn.ec,n runs ni:i(ir in the ciiipty vcssel ant( in t h e vessel
pirtly fillrtl trith short lengths of gla,+ tubing Jrhercby the si~rlaccs1r;ib incwasctl three times.
!'PIE
IN
FIG.I
nm.5
H. ACSTIh- TAYLOR
THE DECOMPOSITIOS O F ETHYLAMISE
TABLE I1 Temperature
szoo ( ' .
Empty Vessel Initial Pressure 129 180 Time mins. Pressure change
Increased Surface 182 mms. Pressure change
12~5
I
21
28
Ii
2
39 48 59
49 68
33 46
82
57
3
1 6 8 IO
76 88 97
15
I11
I
20
I20
25
I2j
I55 162
IO2
i4
11;
86
129
95
46
109 II7.j I22
The close similarity of the rates of pressure change under the two different circumstances suggests a t once that there is little if any heterogeneity in the reaction. The better to demonstrate this absence of heterogeneity the reaction vessel was partly filled with finely divided pyrex glass powder the enormous surface increase thereby involved being expected to show the slightest trace of a surface reaction. The results in Table 111 are typical of many that were obtained. TABLE I11 Temperature j20° C. Empty Vessel Vessel and Powder Initial Pressure 55 5 1 mm Time mms. Pressure Change I
9 5 5
Empty Vessel T esse1 and Powder Initial Pressure j j 51 mms. Time mms Pressure Change 37 5 13
33
IO
5i
15
52
45 50
2
I7
8 15
6 8
3 4
24
22
29
26
41
2766
H. AUSTIS TAYLOR
I t is evident therefore that the reaction is entirely homogeneous. The effects of adding foreign g es were also tried, hydrogen, nitrogen and ammonia being shown to be thout effect. The results in Table I\' of the pressurr change of 5 5 rnms. of ethyhmine and 100 mms. hydrogen are typical of the degree or concordance found with the results in the absence of atitled. diluent.
TABLE IV Time AI?
I
2
3
IO
18
2 j
6 38
4 30
8 43
IO
~j
4;
j2
mins. mms.
Runs were also made in which the pure ethylaniine was nilowed to decompose partially, diluent being then added rapidly and the reaction allowed to eomplete itself. Table 1. gives the results for 60 mms. of ethylamine, I 1 4 mms. of amrnoniR being added a t the end of the third minute. i-
T.IBLI:
Time AP
I
4
2
6
8
IO
15
20
mins.
It ivill be seen that t h e continuity of t h t reiiction is in 110 n.ny disturbed by the aniinonia :idtlition, the values of the pressure increase when plottc(1 against time giving :i smooth curve, eonstzint in direction both hefore nntl :iftilr the ammonia :idtliiion. l l i c decomposition hting thorc3fore undoubtedly homogeneous and unimolecular it should be possible t o calculate tlic velocity constant by mean:: of the usual Corniuln. I t nxiy be seen from the diagrams hoaevcr that in the majority of c:ises, p:irticulnrly n t the 1:ighrr pressures, the total pressure increase a l i e n thc, rc:iction is :lllilnSt complete i j less than the initid pressurc' of ethylaniine taken, for csample 3 9 2 mms. incre:ises at 500' C' by about ,340 itiins. n-lien the rc:iction if almost finis1ied;nt 520' ( ' 2 8 j inms. increases 260 mnie. :ind so on. (-sing t h e gencrnl expression for the velocity constant k of a unimolecular 2
k = __-
io,?
log+a .- x re:iction, n h r r e :I is the initial pressnrc :rnd n - x the difference between the initial pressure and the pressure chnnge after t niinutes, it ip found naturally th;it the values of k fall steadily :is time proceeds. The probable cause of this variation is to b(, found in the polgnierisntion of the eth)-lenc formed during thc clcconiposit ion. I t w n s observed, particularly in t h e experiments a i t h e higher pressures, tlint a few very small droplets of liquid condensed in the capillary tube nttnchetl to the reaction flask :it the point a t which it left the furnace, Tvhen the reaction w:is allnost completetl. Subsequent heating discolorecl this, leaving n brown ring about 2 mms. long a:: the sole evidence t d l tirncls perfectly c h i n . of secondary ticcompositioI1. l'lie flask i t w l f , w 1'e:iw' ant1 I1:iguc. :ind \ \ ~ h r e l c r ?h:ive rcccntly rt ndicd the kinetics of this rc-
t
2767
T H E DECOMPOSITION O F ETHYLAMINE
action. I t has been shown that jo0-550' C is the optimum temperature range for ethylene polymerisation and that above and below this range decomposition into ethane, methane, hydrogen and liquid products is the preferred reaction. The kinetics of the polymerisation are complex, constituting a chain reaction though with probably short chain length. It would seem impossible therefore a t the present time to be able to correct the observed pressure changes recorded above for the amount of polymerisation that may have occurred, particularly since the concentration of ethylene is likewise continually changing. Fortunately, however, the discrepancy between initial pressure and total pressure change is not observed a t lower pressures. The results in the neighborhood of 5omms. show a t all temperatures a strict correspondence. The reasons for this may be many. It is possible that a t these pressures the amount of ethylene polymerisation is actually too small to be observed. Alternatively, the ethylene may decompose producing ethane and a residue which might constitute a pressure increase partly balancing the
TABLE T' Temperature Initial Pressure Time
AP
500
510
55
54
k
AP
520
k
55
k
AI'
-
-
0.5 I
5
7
2
9
I3 18
3 4 6 8
'7
IO
29 34
15
42
23 30 36 40 47
20
25
47 50
52
30
52
53
40
53.5
24
50
Mean Temperature Initial Pressure Temperature
AP -
I
I3
2
21
3 4 6 8
27
32 41 46 47 5
Mean
"352 ,1371 .I3jI ,1362 . I302 .13 18 I329 I35
530
0 5
10
,1389 'I377 I352 ,1388
51
k 2943 . 2 6j q ,2513 ,2468 ,2716
540'c j5 mms. AP k 9.5 17 29 37.5 43 49
'3795 ,3698 ,3747 ,3819 . 3 806 ,3693
,2902
2680 . 2 70
376
9.5 17.5
24
29 37.5 43 47
,1898 ,1916 . I912 ,1874 1909 ,1902 '
52
,1929 ' I940
54
.zoo4
2768
H. AUSTIS TAYLOR
pressure decrease on polymerisation. Finally the polyrrierisat ion may be so slow a t these pressures as to be negligible in comparison with the ethylamine decomposition. I n this connection Pease points out that there is apparently in the ethylene polymerisation a certain minimum initinl pressure below xhich, reaction is extremely slow, a beharior non- considered typical of a chain mechanism. *-It ;oo"C Pease gives this pressure as below j80 mms. I t seems reasonable therefore that the polymerisation of j o mms. of ethylene even a t the highest temperature studied namely 540'C would be extremely small even were it all present during the whole course of the reaction. Actually to test this the pressure decrease of pure ethylene a t j20"C was followed a t various initial pressures. With 2 j o mms. the rate of pressure change amounted to approximately 4 mms. per I O minute interval. At I I O mms. the change was less than 2 mms. in I O minutes whilst a t 40 mms. no change was observed after one hour. I t would seem therefore that the results obtained for the ethylamine decomposition a t the lower pressures would be free from the vitiating conditions obtaining a t the higher pressures. One point may be mentioned a t this stage with regard to the results obtained in presence of hydrogen, where, a t first sight, one might have expected an ethylene hydrogenation to have occurred. Pease shows that the rate of the hydrogenation reaction is of the same order as that of the polymerisation process, ethylene which would have polymerised in absence of hydrogen, appearing as ethane in presence of hydrogen. These results once again however are based on experiments a t a total pressure of about one atmosphere. Calculation of the velocity constants a t pressures around 50 mms. was therefore made a t all temperatures the results being given in Table V. The constancy of the above values would seem to agree with the conclusion drawn above that the data in this range would be free from the variations found a t higher pressures. The possibility of mutually compensating errors in complex reactions must not be overlooked however and the finding of a
TABLE VI Temp.
k
50 0 T jI0
094 I35
j 2 0
192
530
270
5 40
376
E 43,600 43,530 43,240 43,000
l l e a n 43,300 cals. Mean t z , j0o"C 5'0
3 46 2
40
44,080
520
I
io
42,630
530
I
20
5 40
o 86
41,I50 13,?70
Mean 43,500 cds.
2769
THE D E C O M P O S I T I O N O F E T H Y L A Y I N E
constant value is not in itself a very potent argument for its correctness. That the comparative results a t the various temperatures are correct is borne out by the temperature coefficient, I .4for 10' temperature rise, which agrees perfectly with that calculated from the quarter life. The energy of activation calculated from the Arrhenius equation gives values tabulated in Table J?. The data are based on both the velocity constants and the times of 2 5 percent decomposition. The two values are in good agreement and yield a mean energy of activation of 43,400 calories. The accompanying diagram illustrates
,a
,I*
iu
t2r
/a
I z19
l.50
'coo/r FIG.6
the linear relation found between the logarithm of the velocity constant or of the quarter life and the reciprocal of the absolute temperature. The results for the half and three-quarter times are also included for comparison. The agreement of the latter two is not so good since the mean life has been taken a t all pressures. The slopes of the lines in these cases correspond to 40,200 and 38,000 calories for the half and three-quarter times respectively. The results for the data in the 5 0 mms. pressure range when taken alone would obviously give agreement with those calculated from the velocity constants. It is to be observed that the times of quarter or half decomposition show no appreciable change with a change in initial pressure, over the pressure range studied: in other words the reaction shows no tendency to deviate from its unimolecular course within these pressure limits. I n this respect it resembles the decomposition of nitrogen pentoxide' and of azo-isopropane.? There remains still the possibility that the rat? may be lowered at lower pressurrs. It is proposed to investigate this further. Daniels and Johnston: J. Am. Chem. Soc , 43, j3 (1921) See also Ramsperger and Tolman: Proc. Nat. Acad. Sci., 16, 6; Schumacher and Sprenger: 129 (1930). * Ramsperger: loc. cit.
2i;o
H. AUSTIN TAYLOR
For comparison with other unimolecular reactions the value of E R T may be used a t the temperature a t which the rate has a certain value, whether one account for this on a simple kinetic basis as Hinshelwood has done or on the basis of Perrin's' sensibzlzt6 as recently interpreted by D a n i e h 2 Taking the value of the velocity constant as o 00158 sec. -I, the decomposition of KZOSwith an energy of activation of 24,;oo calories gives a t 328' K a value of 38 0 . For acetone the value is higher being 41 4. At 7;3' K the velocity constant of the ethylamine decomposition is o 00157 sec.-* which yields for a n energy of activation of 43,400 a value of E / R T of 28 9. The value is of the same order. Similar values are likewise obtained for other reactions. The decomposition of diethyl ether gives a value of 3 2 9, whilst an extrapolation of the data of Riesenfeld and Schumacher on the unimolecular decomposition of ozone a t low pressures gives the low value of 24.6. Further theoretical discussion is being held over for the completion of work on propylamine now in progress.
Summary The decomposition of ethylamine has been shown to be a homogeneous unimolecular reaction over the temperature range 500-540' C, and a t pressures of 50-400 mms. The energy of activation is 43,400 calories. The value of E / R T of 28.9 is comparable with that of other unimolecular reactions. Nichols Chemical Laboratory New York University, N e w York, N . 1.. 'Ann. Phys., 11 (1919). *Eyring and Daniels: J. Am. Chem. SOC.,52, 1 1 7 2 (1930).
