THEPHOTOLYSIS OF METHYL MERCAPTP~N
Oct., 1 0 6 ~
1431
THE PHOTOLYSIS OF METHYL MERCAPTAN' BY T. INABA~ AND B. DEB.DARWENT Department of Chemistry, The Catholic University of America, Wmhington 17,D . C. Received March 10, 1950
The effects of time, pressure and temperature on the photolysis of methyl mercaptan have been investigated. .4t low conversions hydrogen IS the only important uncondensable gas. Methane arises from secondary reactions involving the product. Hydrogen results exclusively from the breaking of the S-H bond. The decomposition of CHaS requires at least 30 kcal. mole-'. The activation energy and "P" factor of the reaction H CH3SH -c CHIS Hz are about 4.6 f 0.1 kcal. mole-' and 0.34, respectively.
+
+
Introduction The photolysis of methyl mercaptan has been investigated by Skerrett and Thompson3 who found that, on complete photolysis, the gaseous products consisted of HZ(80%) and CH4 (18%) and the condensable products were sulfur and dimethyl disulfide. The quantum yield for the disappearance of CHBSH was about 1.7. They suggested that there may be two concurrent initial photochemical steps CH3BH
+ hv +CH3S + H +CHI +S
(a) (d)
of which the former was the more prevalent and that the hydrogen atom disappeared by H
+ CHISH --ic CH3S + Hz
etry to be a t least 95% pure. Ethylene was a research grade product (Philips Petroleum Company) and was thoroughly degassed before storage. The apparatus was of the type usually used for gas-phase photochemical experiments. A medium pressure Hanovia quartz mercury lamp supplied the active light. Provision was make for mixing the gases thoroughly before photolysis. The temperature was controlled manually to A2O. The volumes of the quartz cell and reaction system were approximately 200 and 300 cm.3, respectively. The gaseous products were separated, analyzed by hot CuO oxidation or mass spectrometrically, and measured by the use of standard techniques.
(1)
Results The first set of experiments was done to investigate the effects of pressure, time and temperature on the rates of formation of Hz and CH4 when the conversion was restricted to less than 0.3%. The data (Table I) show that hydrogen was always the principal product. The formation of CHI varied rather erratically between 0 and 6% of the total uncondensable gas and did not change consistently with temperature, between 50 and 300' or with pressure. In other experiments a t 50' and 54 mm. pressure, where the conversion was extended to
The CH8S radicals combined to form dimethyl disulfide. Although the suggested mechanism is plausible, it is based on rather sketchy evidence. Thus, the experiments of Skerrett and Thompson were carried to extremely high conversion, with the consequence that, a t least toward the end, there might have been serious complications from the photochemical or H atom reactions involving the products. There TABLE I are strong indications4 that the reaction EFFECTO F TIME,PRESSURE AND TEMPERATURE ON THE H
+ (:&OH
+Hz + CHaOH
is important in the photochemistry of methanol, so that the analogous reaction may have been expected with methyl mercaptan. Accordingly, the photolysis of methyl mercaptan has been re-investigated, the conversion being restricted to less than 1%, to obtain further infonnation concerning the nature of the initial photochemical and free radical reactions. Information concerning the rate constants of the elementary reactions has been obtained and an attempt was made to investigate the stability of the free radical intermediates. Experimental The methyl mercaptan (Mathieson Company, 99% pure) was subjected to thorough de-gassing, by trap-to-trap distillations in vacuo, and stored in flasks isolated by mercury cut-offs. Deuteromercaptan (CH3SD) was prepared by repeated equilibration with 99.6% D20 (Stuart Oxygen Company) arid was shown by infrared6 and mass spectrom, (1) This research waa supported by a grant from the Petroleum Research Fund administered by the American Chemical Society. Grateful acknowledgment is hereby made to the donor of said fund. (2) On leave from the University of Osaka, Japan. (3) N P. Skerrett and H. W. Thompson, Trans. Faraday Soc., ST, 81 (1951) (4) AI. K. Phibbs and B. deB. Darwent, J . Chem. Phys., 18, 495 (1950) (5) We are gIatefiil to Dr E. D. Becker and Dr. N. E. Sharpless of the National Institutes of Health Bethesda, Maryland, for the infrared analyses.
PHOTOLYSIS OF CHBSH Reaction cell = 200 cm.8; reaction volume = 300 ~ m . ~ Pressure,b mm.
21 21 115 115 209 209
Time, min.
(a) T 90 180 15 90 15 60
Rates5 HI CHi
= 50' 7.3 7.6 36.9 33.0 43.7 48.0
0.4 0.1 2.3 1.1 2.3 1.1
Pres8ure.b mm.
28 196 196
Time, min.
Ratesa H P CHc
(b) T = 100' 60 8.4 15 40.9 30 39.4
0.4 1.1 0.0
T = 200' 60 13.6 0 . 3 (d) 2' = 300' 60 19.4 .4 30 18.2 .O 22 30 12.9 0 . 0 15 2 7 . 2 1 . 7 151 30 21.3 .O 30 27.7 0 . 9 24C 90 4.3 .O 'igC 90 0 . 8 0 . 0 15lC 90 12.9 .O Rates in cm.3 (N.T.P.) min.-' X 106. Rates of photochemical reactions were obtained by subtracting the thermal rate from the total rate. b Pressures-all corrected to 50". Denotes thermal reaction. (c)
35 78 78 148 148
5.8%, methane became increasingly important; thus, a t 0.61 and 5.8% conversion methane represented 7.5 and 12.8y0,respectively, of the uncondensable gas. In later experiments, with CHJSD, carried to still higher conversions, methane ac-
T. INABAAND T3. DEB.DARWENT
1432
EFFECTOF ETHYLENE ON Pl,
mm.
R,
PZ/PI
(a)
30 30 45 60 60 GO a
4.0 5.0 5.0 2.0 3.0
T
=
6.42 6.41 9.10 11.65 11.01 !0.20
4.0
R, 50 ' 1.90 1.68 2.33 5.21 3.91 3 05
PI,
kdki
mm.
TABLEI1 RATEOF FORMATION OF HYDROGEN"
THE
R.
pp/pI
T
(c)
30 45 60 60
0.60 .57 .58 .61 .61 .58
100
100
-
5.0 5.0 1.0 3.0 2.0 4.0
Ro
3.82 L.1.15 5.56 1.73 7.03 4.85 6.46 2.64 0.53 4.91 10.07 3.79
+ *
+H + Hz
(8)
(1)
are the only important processes in the photolysis. The othcr possible processes
+ hv +CHzSII + H * CHa f HS --+ CHI + S r-r -I- CH~SI-I+CH,SH + II?
CHsSIl
-_-_ (31 \ \ , 1
-\mtic,
iinpubliiherl results.
30 30 30 45
0.46 .44 .44 .48 .47
p2/p1
R.
(b) T = 4.8 4.99 5.0 4.96 5.0 4.52 5 . 0 6.60
R,
kr/kl
120' 1.47 1.46 1.37 1.87
0.50 .48 .51 .51
-
Av. 0.50
.41
-
Av. 0.45
counted for as much as 46% of the uncondensable gas. In those experiments, a t pressure between 90 and 147 mm. and temperatures between 40 and 220°, the uncondensable gas was shown, by mass spectrometric analysis, to be entirely D2and CHID, no Hz or H D was ever detected. The effect of ethylene on the rate of formation of Hz frorn CH3SH was investigated a t a variety of pressures and temperatures. The data (Table 11) show clearly that the rate of formation of H2 decreases with increase in the ratio C2H4/CH3SH. E'urthermore, if R, and R, are the ratesof formation of Hz in the absence and presence of C2H4, respectively, the ratio RaIRp is found to be a linear function of [C2H4]/[CH3SH]and to be independent of the total pressure a t constant temperature. The presence of e1,hylene did not alter the rate of formation of methane, which was always very small. Other experiments, with C3F6 in place of C2H4 show that C:IF6 was also efficient in inhibiting the formation of H2. However, the reproducibility of those cxperirnents was very poor, and there are indications3 that C& changes slowly on storage, possibly by a polymerization or disproportionation type of process. Discussion The results of the present investigation indicate that hydrogen is the only uncondensable product when the conversion is kept small. Methane becomes of significance only when the extent of the converbion of mercaptan is large. Hence i t is likely that CH, does not appear as a result of the photolysis of methyl mercaptan but comes from some pi*ocessinvolving one of the stable products of the photolysis. Methane is not produced early in the photolysis eren at 300°, where the thermal decomposition is becoming noticeable, and the rate of formation of hydrogen is essentially independent of temperature. The results obtained with CHBSDand the fact that olefins inhibit the production of hydrogen show clearly that the reactions
+
Pl, mm.
ka/kr
= 220°
Av. 0.59 Ratchs in cn31.3 min.-' (N.T.P.) X 106.
CHsSH h~ CHB I3 CIIaSII .--) CHIS
Vol. 64
(b) (c)
(d) (1')
do not occur to any appreciable extent. The-fact that, even a t 300' where the thermal decomposition is noticeable, there is no increase in the rates of formation of either Hz or CH4 shows that the CH& radical possesses considerable thermal stability. The process CHsOCHz 4CHzO
+ CHI
requires approximately 19 kcal. mole-' and is easily detectables a t 150'. Hence, if we assume that a t 300' CHZS decomposed a t l/lOth the rate of CHlOCH2 a t 150' and that the frequency factors are the same, the activation energy for the thermal decomposition would be a t least 30 kcal. mole-'. Some information concerning the rate and activation energy of reaction 1 has been obtained from the effect of ethylene on the rate of formation of H2. The treatment of the data is similar to that used by Darwent and Roberts' for HoS. Thus, in the presence of ethylene, the H atom has an alternative method of disappearing H
+ CzH4 *C2H6
(2)
Provided the CzH5disappears without generating H or H2, or by reaction with H atoms, the following relationship holds
. Ra -
1 =4 k xp_2 RP ki PI where pl = [CHISHI and p2 = [C2H4]. Hence, by measuring R, and Rp a t known p , and p2 we can calculate the ratio kz/kl. The above relationship requires Ra/R, to be a linear function of p2/pl, and independent of the total pressure. The results show that those requirements have been met. In particular, the fact that the value of k2/kl is not sensitive to changes in the total pressure is strong indication (a) that the H atom formed in reaction (a) does not possess energy significantly greater than the thermal equilibrium value or, if it does, it is nearly always deactivated by collision before it reacts, and (b) that the CzHs formed in reaction 2 does not, under the conditions of the experiments, redissociate into CzHa and H. The effect of temperature on k z / k l has been examined from an Arrhenius plot of the results of individual experiments. Although the points are rather scattered, the extreme values of AE = E1 E2 are 0.42 and 0.67 kcal. mole-'. Adopting the average value of 0.54 f 0.12 for AE and assuming (6) R. A . Marcus, B. deB. Darwent and E. W. R. Steaeie, J . Chem. Phya., 16, 987 (1948). (7) B. deB. Darwent and R. Roberta. Faradolj Soc. D i w . . 14. 55 (19.53).
Oct., 1960
I < I N E T I C S O F T H E RC.4CTION O F
AROMATICI I Y n R O C A R R O N S
Z1 = ZZ,we find P2/P1 = 0.26. If we accept the previously' suggested valuesof EZ= 4.1 kcal. mole-' and Pz = 0.09, me find E1 = 4.6 kcal. and PI = 0.34. For the reaction of D with DzS, the previously' derived values were E = 5.0 kcal. mole-' and
IN
SULFURIC I!CIT)
1433
P = 0.73. The slightly lower activation energy and P factor for reaction 1 is not unexpected since substitution of H by CHs usually results in a lower dissociation energy of the remaining bond and also halves the intrinsic probability of the reaction.
TIIF, KIK'ETICS OF THE REACTIOSS OF AROMATIC HYDROCARHONS I N SULFURIC AC1D.l I. BESZESE BY MARTINKILPATRICK, MAXW. MEYERAND MARYL. KILPATRICK The Department of Chemistry, Illinois Institute of Technology, Chicago,Ill. Received March 18. 1900
The rate of sulfonation of benzene in sulfuric acid solution has heen measured spectrophotometrically; the results ohtained are in fair agreement with those of Gold and Satchell, obtained by an isotope-dilution technique. An empirical equation has been found which fits the experimental results and which can be interpreted in terms of the concentration of the molecular sulfuric axid present in mixtures of sulfuric acid and water, as determined by Young and co-workers using Raman spectra, and the activity of water in these mixtures, as determined by Giauque and co-workers.
In an attempt to unravel the kinetics of the Jacobsen reactic1nl2it became necessary to understand the kinetics of sulfonation of the four- and five-methylsubstituted benzenes. The range of concentration of sulfuric acid for pure sulfonation is limited for durene (1,2,4,5-methylbenzene), isodurene (1,2,3,5-methylbenzene) and pentamethylbenzene by desulfonation, isomerization and disproportionation reactions. To avoid complications, the first series of experiments was carried out with benzene. The rate of reaction was determined by following the appearance with time of benzenesulfonic acid spectrophotometrically in the ultraviolet. A Cary Spectrophotometer Model 11 equipped with a program attachment was used. The reaction vessel was the absorption cell, and the general procedure was as follows. Sulfuric acid of known concentration was cooled in a refrigerator below the temperature planned for the experiment, a drop of hydrocarbon added from a syringe 1.0 give a concentration of 1 X 10-4 mole per liter, and a.fter shaking for five minut'es the absorption cell was filled and placed in the thermostated compartment of the spectrophotometer. After allowing a suitable time for the establishment of thermal equilibrium, the recording spectrophotometer was started and the rate of formation of sulfonate determined a t the absorption peak 272 mp. The molar extinction coefficient for the benzenesulfonic acid is so much greater than that for benzene that no appreciable error is introduced by the presence of unreact,ed hydrocarbon. Experiments were carried out at 25 and 12.3"with t,emperature control of =k0.05" in most cases. As the r e a d o n proved to be strictly first order in hydrocarbon, as shown by a typical run in Fig. 1, in most later experiments the Guggenheim3 method of evaluating the velocity constant was used. The quc'ted rate constants have limits of (1) Presented in part a t the 134th Meeting of the American Chemical Society, Chicago. #jePtemher1958, before the Division of Physical Chemistry. (2) 0. Jacobsen. Bet., 19, 1209 (1886); 20, 896 (1887). (:;) E. A . Ciimzetilicim, Z'hil. M a g . , 1 , 5.38 (19%).
error of &5% (as compared with +20% in the work of Gold and Satchell). The equation d[ArS03H1 = ko~,s
