THE THERMAL DECOMPOSITION 0 F DIMETHYL SULFITE1 RICHARII C WOOD
ATD
H AUSTIS TAYLOR
Department ot" Chernzstrij, -Vew I'or,k ~ n i c e r s $ ? j ,New Y o r k , N e w York Received March 26, 19bY
The thermal decompositions of organic compounds studied, especially during the past ten years, in the search for simple homogeneous unimolecular reactions have added considerably t o our knowledge of the mechanism of numerous pyrolyses. The enormously different behavior found in the decomposition of methyl nitrite (6) and of its isomer nitromethane ('i)-the former essentially simple, the latter extraordinarily complex-is but one illustration of the variety which confronts the worker in this field, especially where little is known of the bond strengths in the compounds used. A search of the literature on organic sulfites shows but one reference to their thermal stability or to the products formed on pyrolysis. Prinz (4)states that diethyl sulfite boils with decomposition a t 2OO0C., giving sulfur dioxide and ethyl ether. It was on the basis of such a possible simplicity t h a t this investigation \vas undertaken. The choice of the methyl compound was made solely for its lower boiling point and with the knowledge t h a t methyl ether, if formed, should be relatively stable a t such low temperatures. .is is later shown, however, the reaction is by no ineans so simple, a t any rate for the methyl derivative, and there appears to be a distinct possibility that Prinz may actually have been working with the isomer, ethyl ethylsulfonate. The dimethyl sulfite was prepared by the method of Arbusow (1). Two moles of freshly distilled thionyl chloride was allowed t o drip slowly into four moles of dry methyl alcohol. During the early stages the mixture must be cooled below 20°C., othern ise the yield of sulfite is negligible. Hydrogen chloride mas carried off in a stream of nitrogen bubbling through the reaction mixture, which was thus kept effectively stirred. After the first violent reaction subsided the mixture was allowed to warm up and was finally refluxed on a water bath for two hours t o remove all hydrogen chloride. Fractionation of the product yielded a sample boiling between 1 Abstract of a thesis submitted by Richard C. Wood in partial fulfillment of the requirements for the degree of Doctor of Philosophy a t S e w York University, June, 1936.
1091
1092
RICHARD C. W O O D AND H. A U S T I S TAYLOR
125.8 and 126.1"C. It has a n odor very similar to t h a t of acetone, without a n y trace of the sharpness characteristic of sulfur dioxide. Samples of the sulfite were allowed t o stand in glass-stoppered bottles for some months during the course of the work, but on distillation gave no evidence of sulfur dioxide and showed no change in boiling point. To determine the general course of the pyrolysis and the products of decomposition, the sulfite vapor was passed through a Pyrex tube maintained a t 550°C. The temperature used resulted from experience of the extent of decomposition occurring during the very short time of contact. The vapors passed through three traps on issuing from the furnace, the first a t room temperature, the second in ice, and the third in a dry icetoluene mixture. All gases escapiqg beyond this were collected over water. The products in the first trap proved, on distillation, to consist primarily of unchanged sulfite, together with a lower boiling fraction with a very disagreeable odor markedly resembling, by comparison,, t h a t of dimethyl disulfide. The quantity did not permit a n accurate determination of the
SAMPLE
1
3"
1
1 I
,
___
TABLE 1 Analysis of the exit gases so2
I
Hz
j
CHI
j
35 .O 27.2 27.8
__ .-
24 2 26.8 27.4 32 8 ~- -
22 4 22.0 21 .-I 29 2 ~
13.6 17.0 16 6 -
___-_
boiling point, which, however, was less t h a n 110°C. The residue from this distillation on standing yielded white needles similar t o those observed on the cooler parts of the pyrolysis tube just outside the furnace. After recrystallization from alcohol, a mixed melting point determination a t 109°C. proved them to be dimethyl sulfone. The second trap (in ice) was empty. The third tyap (at -78°C.) contained a quantity of liquid which analysis showed to be sulfur dioxide. Specific tests showed the absence of sulfur trioxide, mercaptans, and hydrogen sulfide. A small amount of a yellow solid was observed in the connecting tubes between the traps. It did not volatilize a t room temperature but was not sulfur. It proved to be highly unstable and in the presence of dilute nitric acid gave off copious quantities of sulfur dioxide. The quantity being small, further identification was not attempted. The gases collected over water were analyzed in a Fisher gas analyzer. The absence of oxygen, carbon dioxide, and unsaturated hydrocarbons, and the presence of carbon monoxide, hydrogen, and methane were first
THERMAL DECOMPOSITIOS O F DIMETHYL S U L F I T E
1093
demonstrated. With a view to determining the relative amounts of these gases in comparison with the sulfur dioxide formed, a complete analysis was made of the exit gases without freezing out the sulfur dioxide. The analysw in table 1 were found. If the decomposition were to be represented as:
(CH3)zSOS + SOz
+ CO + CH, + Ha
the analyses given in table 1 would prove to be somewhat lorn in carbon monoxide and still lower in hydrogen. The presence of considerable amount. of diincthyl sulfone would involve a reduction as follows:
which might account for the hydrogen lacking in the analysis. N o esplanation of the low carboii monoxide values can be offered other than incomplete absorption in thc analyses, where it will be observed t h a t the totals also are less than 100 per cent. Having thus obtained some knowledge of the products of the decomposition, measurements of the rates under varying conditions of temperature and pressure were attempted by a static method. The reaction vessel was a 236-cc. Pyrex flask mounted in an electric furnace, manually operated, thc temperature of which, measured by a platinum resistance thermometer, could be controlled to half a degree. The flask was connected by capillary tubing to a mercury manometer, a storage supply of sulfite, and a Hyvac pump. The connecting capillaries were heated to a temperature around 130°C. to prevent condensation during an experiment and t o maintain a reasonably high vapor pressure of the sulfite. This involved also maintaining one stopcock at this temperature. A stopcock lubricant with a Dammar gum base was found satisfactory. The temperature range covered was from 340 to 390°C. At 340°C. and 35OoC.,where the reaction rate was quite slow, only a search for possible evidence of an induction pfriod in the reaction was made. N o such evidence wau found even at low pressures. The pressures of sulfite initially present covered the rangc from 10 to 250 1mn. I n vien- of the large pressure increase the latter figure was the practical upper limit 11-ith the apparatus employed. It was found that the rates of reaction on a clean surface were always high and erratic. When the reaction \-esse1 had been evacuated for from fire to fifteen minutes after a run and a second run then made, the rate of reaction was lon-ered and could bc more easily reproduced. Such a pumping procedure was standardized for all runs, data for which are given in table 2. The products of the reaction appeared t o be without effect on the rate, but undoubtedly a clean surface is poisoned during a run although, as will be seen later, not sufficiently t o suppress completely a heterogeneous reaction. THE J o r R m L OF PHYSICAL CHEXISTRS,
YOL. 41, s o . 8
I094
RICHARD C . WOOD A S D H. IIUST1I-i T d Y L O R
Data of a typical experinleiit are given in table 2. An attempt to calculate unimolecular velocity constants from the rate of pressure change directly observed, gave values which rapidly decreased during a single run. The constants given in table 3 are therefore the ralues a t zero time extrapolated from readings a t 15-second intervals over the first two to five minutes of reaction. Included in table 3 are data obtained using the reaction vessel packed with short lengths of Pyrex tubing with fire-polished ends, whereby the surface to volume ratio wsls increased about elevenfold. The rate is seen to be increased by a factor of two, indicating definitely some heterogeneity. If the homogeneous and heterogeneous reactions are independent of each other and the calculated elevenfold increase in surface may be relied upon, this doubling in the rate should correspond t o about 11 per cent of the pr2mLri 2 __
Data u f c t ~ p z c a le z p e r ~ m ~ n t Temperature, 370°C ; initial pressure, 1-10 min.
______
-
- -~
I
TIME
PRESSURE IVCREhSE.
__
.~
minutea I
1
2 3
5
10 13 20
I
PRESSURE INCREAPE
TIME
-- __ - -
?I1 I l l .
mtnzltes
19 37 52 78 100 128 165 196
25 30 40 50 60 80 100 120
I
I I
~
1
I
mm.
220 238 266 287 303 328 310 344
1 ___ -~
--
total reaction being heterogeneous. -1study of the actual rate of pressure change on the increased surface showed that the rate falls off faster than the rate on the original surface. The increased surface, in other words, causes a marked acceleration only in the early stages. This might seem t o suggest a progressive poisoning of the increased surface. The results, however, were quite reproducible, and more prolonged seasoning was without effect. Furthermore, the percentage pressure changeon the increased surface was slightly less than on the original surface. This latter value, however, was variable with temperature, increasing with increasing temperature from 240 t o 260 per cent a t temperatures from 360 t o 390°C. These values are not too reliable, since on very long standing a small pressure decrease, sometimes amounting to as much as 5 per cent, was obserred. This was undoubtedly due to a diffusion of some of the products from the reaction vessel and condensation on the cooler connecting tubes. The reaction vessel a t all times was perfectly clean, but a black deposit
1095
THERMAL DECOMPOSITIOS O F DIMETHYL S U L F I T E
collected in the capillaries which could be distilled along them leaving a clean surface behind. The above data are plotted in figure 1 as log KO against log P . Over the pressure range covered the points lie on straight lines, showing a n increase in K Oas the pressure increases. The slopes of the lines are the TABLE 3 Coinplete decomposition data in velocity constants at zero time I
TEYPERATURE, 3 6 2 ° C
TEMPERATURE,
I
-~
- -
K~ x
Initial pressure
I
io8
ni i i i
(e(.
23 24 33 75 96 186 251
5 52 6 15 5 92 6 25 6 62 6.82 7 08
I I I
_______
-1
ni v i .
I
17 32 50 98 140 20 1 257 -
1
1
TEMPER~TURE.380°C.
Initial pressure
___ ___
KO
_
see.
13 22 27 65 120 181 235
13.8 16 1 17 6 17.7 18 6 19 4 21 1
_
_
Initial pressure
~
, IVCREASED
sec. -1
I
8 19 8.97 9 01 9 83 10 2 11 4 11.1
~- -______ -~
-1
39OOC
Initial pressure
102
I n Ill.
TEUPERATLRE, 3 7 0 ° C
_
K~ x 103 I
TEMPER4TURE,
x
370°C.
Initial pressure
~
II
~
1
, ___ _
K OX 10'
mni.
see.?
11 27 28 42 72 121 153
27 9 30.2 30.2 34 8 34 2 37 1 37 1 -
-
_ __ -
-_
STRF~CE
-
_______
-
~
-
__
K~ x 103
IILI??.
bee. -1
23 33 126 165 231
16.9 17.1 20.8 21 3 22.0
I
I
same, namely 0.115, indicating the reaction to have a n overall order slightly higher than the first. It might appear that the reaction was pseudounimolecular, but already in a pressure range where the equilibrium quota of activated molecules was not maintained. What is probably more likely, however, in view of the complexity of the products, is t h a t the reaction
1096
RICHARD C. WOOD AKD H. A U S T I N TAYLOR
proceeds by a free radical mechanism involving several simultaneous reactions. That the reaction is not a chain reaction, or, if a chain reaction, is of very short chain length was demonstrated by attempting to induce sulfite decomposition a t 300°C. by free methyl radicals from azomethane. I n table 4 are given the actual pressure change data for two mixtures of the sulfite and azomethane. It will be seen t h a t the rate of pressure change roughlyparallels t h a t found for azomethane and shows no indication of any decomposition of the sulfite whatsoever, since the pressure increase observed is approximately 100 per cent of the azomethane initially present, whereas had the sulfite
1.6
-
390' 0
--cT
Q 14-
x
c? 0
0.8
-
06
-
--+--I
I
I
I
t
I
I
LOG ?
FIG.1. Changc of initial rate with change in pressure
also decomposed the total pressure would have risen t o well over 1 atmosphere in the second case above. It may be remarked t h a t no particular effort had been made to purify the azomethane, though its ability to induce decomposition in other substances was demonstrated. Since the data plotted in figure 1show parallel lines, the energy of activation is independent of pressure over the pressure range studied. From the temperature coefficient of these lines a n average energy of activation of 55,000 cal. was calculated. The effect of added gases on the initial rate of reaction was determined for nitrogen and for the products of the decomposition. It was found t h a t traces of oxygen in the nitrogen caused a considerable acceleration of the rate of pressure change. On removing these traces, however, no effect for
10%
THERMAL DECOMPOSITION OF DIMETHYL S U L F I T E
nitrogen alone could be found, the results duplicating those found in its absence. The effect of the products was tested by allowing a small amount of the sulfite t o decompose completely and then adding a fresh amount of sulfite whose rate of decomposition was then carefully followed. This was done for varying ratios of sulfite to product a t both high and low total pressures. The results obtained were always in agreement with the data found in the absence of added productswithin the experimental error. Details therefore need not be given.
__
TABLE 4 Decomposition. in the presence of azoniethniLc Temperature, 300°C.
__.~_~._____._~
~~~
. .. .
INITIAL S U L F I T E . . .,. ISITIAL .42OUETHASE.. . .
. . . 3 9 mm. . . . 2 0 mm.
Time
I
Pressure
man.
1
ntm.
4
61 62 64 65 66.5
5
68
0.5 1 2 3
10 15 20
30
40
I
i
72 75 77 79 79.5
__
ISITIAL SULFITE. , , . . . . . . . . ,300mm. I R I T I A L A Z O M E T H I A E . . . , , . 20 mm.
.
'I
Time
I
1
min.
0.5 1 2 3 4 5
, I
I
I I
I'
10
I I
I1
15 20 30 40
Pressure TII~L.
321 322 324 325.5 326.5 327.5 332.5 336 337 339 341
DISCUSSION O F R E S G L T S
The work of Prinz previously referred t o and the major products of the reaction here found would suggest t h a t the initial decomposition is a split of the sulfite into methyl ether and sulfur dioxide, followed by a decomposition of the ether into methane, carbon monoxide, and hydrogen. Against this, however, are the following facts: ( 1 ) t h a t every effort to find methyl ether among the products failed; ( 2 ) that methyl ether alone decomposes a t a reasonable rate around 450 to 550°C. (2); and (3) t h a t the energy of activation of methyl ether, even in this temperature range, is 58,500 cal. (2). Since the energy of activation here found for dimethyl sulfite is only 55,000 cal., the probability of a n initial decomposition into a n already activated methyl ether molecule, which would facilitate its decomposition, seems small, unless the reaction (CR30)2SO -+ CHSOCHZ
+ SO2
1098
RICHARD C . WOOD A S D K. AUSTIS TAYLOR
were fairly highly exothermal. Lacking any absolute data, this would not appear reasonable. On the contrary it might be expected t h a t the reaction was actually endothermal. Reverting to Prinz's statement, it is now known t h a t diethyl sulfite, (C,HSO),SO, boils a t lG;OC., whilst the isomer, the ethyl ester of ethylsulfonic acid, C2H6SO3C2H,boils a t 206°C. It would appear probable, therefore, t h a t Priiiz was working with the latter compound. This is the inore reasonable, too, in view of the two coordinate linkages in the latter compound.
It is hoped that a report on this compound limy shortly be made. The only remaining possibility for a niechanisni of the decomposition of the sulfite would seeni to involve free radicals. It is a relatively simple matter to work out a scheme which will yield the observed products. It is not so simple, however, to see how such schemes should fail t o yield chain reactions of chain length a t least comparable to those suspected (3,5) in methyl ether, unless such a reaction ab
2CH3
+ SO?
(CH3)zMOz
-+
were a very efficient chain-breaking step. S o .pecifie c-tiiiiate m s niadc of the aniouiit of dimethyl sulfone formed during thcl reaction, though the suggested earlier, however, there is quantity was quite conridcrablr. no reason to suppose that the presence of the sulfone may not be due to n reduction of the sulfite. The sulfone boi1.i at 238°C. and shows a marked resistance to both oxidation aiid reduction and might therefore conceivably be stable in the temperature range 360 to 390°C. used here. It seenib probable t h a t this was the product condensing in the capillaries during an experiment and, judging from the end points on the increased surface, may be the product responsible for the small heterogeneity observed. There would appear to be nothing to gain in considering specific free radical schemes, in the abieiice of more information. It is anticipated t h a t an investigation of the sulfonate isomer niny ihed more light on the complexities here found in the sulfite. SUMMARY
The decomposition of dimethyl sulfite, studied in the temperature range 360 to 390°C. and a t pressures from 10 to 250 min., yields methane, carbon monoxide, sulfur dioxide, hydrogen, aiid dimethyl sulfone as principal
THERMAL DECOMPOSITIOS O F DIMETHYL S U L F I T E
1099
products by a reaction of order slightly higher than the first, which is not, however, entirely homogeneous. It is presumed that the reaction involves free radicals, but it is shown that if it is a chain reaction, it has only short chains. The overall energy of activation is 55,000 cal. REFERESCES (1) ARBUSOTV: Cliein. Zentr. 11, 885 (1909). (2) HISSHELWOOD A ~ V DASKEY:P r o c . Roy. SOC.(London) A116, 215 (1927). (3) LEERMAKERS: J. Ani. Chem. S O C . 56, 1899 (1934). (4) PRIXZ: h n . 223, 374 (1884). ( 5 ) RICEAND HERZFELD: J. Am. Chem. SOC.56, 284 (1034). (6) STEACIE ASD SHIW:Proc. Roy. Soc. (London) A146, 388 (1934). (7) T.LYLOR ~ S DVES,ELOVGKT: J. Phys. Chem 39, 78 (1936).
