C. E. WARINGAND REMOPELLIN
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The Kinetics and Mechanism of the Thermal Decomposition of Dimethylmercury
by C. E. Waring and Rem0 Pellin Department of Chemistry, University of Connecticut, Stows, Connecticut
06868 (Received October 20, 1966)
The kinetics of the pyrolysis of dimethylmercury have been investigated over a pressure range from 2.0 to 90.3 mm between 312 and 401.1'. The reaction is homogeneous and first order in a conditioned reaction flask. The predominant gaseous products of the decomposition are methane and ethane. Nitric oxide produces inhibition with the result that the rate of disappearance of DMM is reduced and the ethane concentration falls almost to zero. Propylene also inhibits the pyrolysis but hydrogen catalyzes it. A reaction mechanism is postulated which is in accord with the experimental results.
Introduction
A recent paper by Kallend and Purnell' has pointed out the lack of agreement in the experimental results obtained by various investigators2on the pyrolysis and photolysis of dimethylmercury (DMM). Here, again, the results and interpretations of these authors are not wholly in agreement with the previous papers in certain important areas. Our study of the thermal decomposition of DMM was actually undertaken with the aim of determining the role of nitric oxide in the presence of methyl radicals. This compound was chosen because it appeared to be a simple and uncomplicated source of methyl radicals. In order to better understand the reactions of NO, propylene, and other foreign gases, it was, of course, first necessary to investigate thoroughly the kinetics and mechanism of the decomposition of DMM. The data presented here extend the work of Kallend and Purnell. I n some instances our results are in agreement with theirs; in others, there appears to be no ready explanation for the wide disparity in the experimental observations. Although many of the data presented here were obtained prior to the report of Kallend and Purnell, a number of experiments have been subsequently r e peated and extended. It was thought worthwhile, therefore, to present our results in the hopes that the similarities and differences between these data and those of others may eventually lead to a complete unThe J O U Tof ~Physical Chemistry
derstanding of this reaction and the effects of foreign gases on it.
Experimental Section A . Apparatus. The rate studies were made using a static manometric system. The reaction chambers were cylindrical Pyrex tubes of 150-ml capacity. Before using, each was uniformly coated inside by a layer of carbon. The reaction flask was placed in a copper tube which in turn fitted into the well of an electric furnace. The temperature of the furnace was controlled by an electronic regulator to within k0.2" as measured by a calibrated nitrogen-filled mercury thermometer. The temperature inside the copper tube was found to be constant to within k0.05" over its middle 12 in. The maximum temperature variation of the furnace over a 24-hr period was k0.5'. Because of the corrosive action of dimethylmercury on stopcock lubricants, mercury cutoff valves were employed. To prevent condensation, the vacuum (1) A. S. Kallend and J. H. Purnell, Trans. Faraday Soc., 60, 93, 103 (1964). (2) H. W. Thompson and M. Meissner, Nature, 139, 1018 (1937); J. P. Cunningham and H. 9. Taylor, J . Chem. Phys., 6 , 359 (1938); J. H.Raley, F. F. Rust, and W. E. Vaughan, J . Am. Chem. Soc., 70,88 (1948); F. P. Lossing and A. W. Tickner, J. Chem. Phys., 2 0 , 907 (1952); K.U. Ingold and F. P. Lossing, ibid., 21,368, 1135 (1953);
L. M. Yeddanapalli, R. Srinivasan, and B. Paul, J. Sci. I n d . Res. (India). B12. 232 (1954): C. M. Laurie and L. H. Long. Trans. hararadab i o c . ; 51, 665 (1955); J. Cattanach and L. H. Long, ibid., 56, 1286 (1960).
THEKINETICSAND MECHANISM OF
THE
2045
THERMAL DECOMPOSITION OF DIMETHYLMERCURY
line was wrapped in Nichrome ribbon and heated electrically to a controlled 75'. The reaction system was evacuated by the usual techniques and before each run the reaction system was evacuated until the pressure was mm or less as measured by a McLeod gage. This, together with the flaming out of the vacuum line before each run, eliminated the possibility of oxygen contamination. The gaseous reaction products were removed for analysis by means of a Toepler pump which was attached directly into the line. The analyses were made by a modified Bone-Wheeler technique, a vapor phase chromatograph, and by a mass spectrometer. B. Material. Dimethylmercury was prepared by the method of Marvel and GouldS3 The physical constants for this preparation were bp (cor) 94.8', N d (22.2') = 1.5325, and density (22.2') = 2.9542. These values were in good agreement with those in the literature4 of 95', 1.5327, and 2.95412, respectively. Nitric oxide was prepared by the method of Johnston and Giauque.6 The hydrogen, nitrogen, helium, and propylene employed were Matheson Co. CP grade. The hydrogen was passed over copper oxide at 400' and through a Dry Ice trap before use. Before the propylene was introduced into the reaction chamber, it was first bubbled through dibutyl phthalate. The other gases were used without further purification.
Results 1. Nature of the Decomposition. DMM was found to decompose at a conveniently measurable rate a t temperatures between 312.0 and 401.0' and a t initial pressures from 2.0 to 90.3 mm to give predominantly methane and ethane. The over-all profile of the decomposition as observed by pressure-time readings is unusual, as is seen in Figure 1. The graph shows that the pressure first increases to a maximum value, p,, which is nearly twice that of the initial pressure, pi. The pressure then decreases to a final pressure, pt, which is about 40% greater than the pi. At this point, no further decrease in pressure with time occurs. This phenomenon was observed at all eight temperatures investigated and was readily reproducible. The formation and disappearance of this maximum is undoubtedly due to two concurrent effects. It was observed that finely dispersed mercury droplets condensed on the walls of the tubing extending from the furnace during the phase of the decomposition when the pressure was decreasing from the maximum. The other effect was a yellow translucent film which also deposited on the exit tube outside the reaction flask when the gaseous products were being evacuated.
80
8 60 X R
';i 40 4
20
0
50
100
150
200
250
Time, min.
Figure 1. Pressure change with time for dimethylmercury ( p i = 36.0 mm) a t 8, 370.0'; 0 , 384.8"; and 0,401.0'.
It seems evident, therefore, that the decomposition of DMM is complicated by the simultaneous occurrence of a condensation and polymerization process. It should be noted in Figure 1 that at 401', P, is reached a t approximately 20 min reaction time. Table I presents the variation of the methane and ethane concentration with time a t 401 ', Table I : Partial Pressures and Mole Per Cents of CH, and CzHsfrom the Decomposition of 30 mm of DMM a t 401.0' Time, 8eo
PCli49
PCaEsv
%
%
mm
mm
CH4
CzHa
0 40 60 120 240 600 840 3,300 15,000
0 5.0 7.4 10.6 14.4 16.7 17.6 18.0 18.4
0 4.0 5.7 7.6 10.4 12.2 13.0 13.2 13.0
0 55.6 56.5 58.3 58.1 57.8 57.5 57.7 58.6
0 44.4 43.5 41.7 41.9 42.2 42.5 42.3 41.4
These data clearly show that the decomposition of DMM is essentially complete a t the pressure maximum. Consequently, insofar as the formation of methane and ethane are concerned, p , may be taken to be equivalent to the final pressure of this reaction. The subsequent decrease in pressure, then, refers to the condensation and polymerization processes which (3) C. 5. Marvel and V. L. Gould, J. Am. Chem. SOC.,44, 153 (1922). (4) "International Critical Tables," McGraw-Hill Book Co., Inc., New York, N. Y., 1933. (5) H. L. Johnston and W. F. Giauque, J. Am. Chem. SOC.,51, 1394 (1929).
Volume 71, Number 7
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C. E. WARINGAND REMOPELLIN
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become predominant after the decomposition is complete. Figure 2 shows that the ratios of p,/pi and pr/pi are invariant with temperature and with initial pressures between 20 and 90.3 mm. This constancy lends additional support to the contention that the decomposition of DMM proceeds independently of the subsequent processes. The slight increase in these ratios at the lower initial pressures is probably due to the decomposition of condensation products which were found to occur at the higher pressures. 2. Order of Reaction. Table I1 presents typical data for the change of pressure with time for various initial pressures.
Table 11: Change in Pressure (mm) with Time for Various Initial Pressures of DMM a t 401' P I , mm
1, aec
6.0
0 3 5 8 10 20 30 60 90 120 150 180 210 240 270 300 360 420 480 540 600
0 0.12 0.20 0.32 0.40 0.85 1.27 2.05 2.62 3.18 3.50 3.!)0 4.14 4.42 4.56 4.86 5.14 5.36 5.58 5.76 5.88
14.5
0
0.30 0.49 0.79 1.00 2.10 3.10 4.98 6.38 7.64 8.48 9.39 10.03 10.68 11.23 11.81 12.60 13.19 13.70 13.99 14.28
23.5
34.7
48.0
70.2
90.3
0 0.54 0.84 1.33 1.66 3.33 5.02 8.34 10.78 12.64 13.81 15.32 16.45 17.39 18.35 19.17 20.18 21.05 21.85 22.32 27.70
0 0.72 1.21 1.97 2.39 4.85 7.14 12.11 15.92 18.21 20.29 22.52 24.46 25.67 27.10 28.28 29.84 30.88 31.92 32.79 33.13
0 0.96 1.63 2.64 3.26 6.62 9.69 16.80 21.88 25.92 28.32 31.58 33.84 35.61 37.44 39.31 41.71 43.68 44.88 45.60 45.88
0 1.54 2.38 3.86 4.70 9.68 15.16 24.99 31.66 38.04 41.41 45.77 49.14 51.94 54.47 57.07 60.23 62.47 64.58 65.28 65.98
0 1.89 2.97 4.87 6.05 12.37 18.96 31.51 41.17 48.40 52.82 58.33 62.93 66.46 69.71 73.14 77.38 80.36 83.25 84.43 84.88
2.1
-
0
f
e
-
-
1.8
a
\
1.2
0.9 0
20
40 60 p(initial), mm.
80
100
Figure 2. Ratio of the maximum pressures (upper curve) and final pressures (lower curve) to various initial pressures of DMM at 401.0' (o),384.8' (e),and 370.0' (0).
partial pressures of methane and ethane for various initial pressures of DMM at 401 O . These data were also programmed on the IBM 1620 computer to solve for the initial rate constant, ki, and n, the order of reaction for both methane and ethane in eq 1 . Values of n = 0.990 and 1.00 were obtained for methane and ethane, respectively, again indicating first-order kinetics. 3. Energy of Activation. A plot of the logarithms of the initial rate constants, ki, us. the reciprocals of the absolute temperature at eight different temperatures gave an excellent straight line, as seen in Figure 4. The activation energy was calculated by the method of least squares, and the rate constant for the thermal decomposition of DMM, expressed in terms of the Arrhenius equation, is
hi = 9.04 X 10ge-37,300'RT sec-1
(2)
These data were programmed on an IBM 1620 computer to solve for ki, the initial rate constant, and n, the order of reaction, in the equation log (dp/dt)o
=:
log ki
+ n log pi
(1)
A value of n = 1.01 was obtained, indicating that the order of the decomposition is unquestionably first. The linearity of the curve was further checked by computing the second-order term, which proved to be negligible. Visual evidence for the firsborder kinetics of this reaction is given in Figure 3. Table I11 presents typical data for the variation of The Journal of Phusical Chemistry
I 0
120
I 240 Time, sac.
I 360
I 480
I 600
Figure 3. Pressure change with time for dimethyl mercury at 401" at initial pressures: 0,6.0 mm; 0, 14.0 mm; e, 23.5 mm; (3, 34.7 mm; 0, 48.0 mm; 0 , 70.2 mm; 0,90.3 mm.
THEKINETICSAND MECHANISM OF
THE
THERMAL DECOMPOSITION OF DIMETHYLMERCURY
2047
Table 111: Variation of Partial Pressures of CH, and CnH6 with Time for Different Initial Pressures of DMM a t 401’ -14.5-CH4
1, sec
0 1.9 3.5 5.6 6.5 7.0 7.5 7.8 8.1
0 30 60 120 180 240 360 480 600
CzHe
-23.5CHI
CzHe
CHI
CzHs
0 3.2 5.7 8.9 10.5 11.3 12.2 12.6 13.1
0 2.3 4.1 6.5 7.6 8.1 8.9 9.4 9.9
0 4.1 7.4 10.6 13.4 14.4 15.6 16.1 16.7
0 2.9 5.7 7.6 9.4 10.4 11.4 12.0 12.2
0 1.4 2.4 3.8 4.5 5.0 5.5 5.8 5.9
-30.-
3.5
7-48.CH4
0 6.6 11.8 18.4 21.4 23.1 24.9 25.8 26.7
F
I
1.45
1.50
CzHs
-70.2-CHI
0 4.6 8.3 12.8 15.0 16.7 18.3 19.2 19.5
0 9.6 17.8 26.8 31.3 33.6 36.4 38.0 39.2
1.55
1.60
CzHs
0 8.2 13.3 19.1 22.3 24.2 26.6 27.9 28.5
3.0
,. P
2.5
5 2.0 3
2 1.5
4
1.0
1.450
1.500
1.550 1.600 lO*/T, OK-*.
1.650
1.700 1.725 1.65
10*/T, OK-’.
Figure 4. Determination of the activation energy for the thermal decomposition of dimethylmercury.
Figure 5. Determination of the activation energies for the formation of methane (0) and ethane ( 0 )during the decomposition of 30 mm of DMM.
The activation energies for the decomposition of
DMM as measured by the initial rate of formation of methane and ethane were also obtained. In these cases, the rate constants, kCH4 and kC2Ha, were found by dividing the number of moles of each gas at various times, t, by the time and by the original number of moles of reactant. Table I V gives the rate constants obtained at four different temperatures and Figure 5 presents the plots of the logarithms of these rate constants against the reciprocals of the absolute temperatures. Table IV: Variation of the Rate Constants of CHa and C2Hawith Temperature
x
TO
kCHi 10’ sec -1
350.5 370.0 384.8 401 .O
0.621 1.22 2.36 4.30
kCnHs x sec -1
10’
0.215 0.538 1.45 3.36
Again, the activation energies for the formation of these two products were obtained by the method of least squares, and the rate equations are
kCCH, = 6.84 X 109e-32,200/RT sec-’ kCzHs
= 1.93 X 10ize-45”00’RT sec-‘
(3) (4)
If these energies are weighted according to the number of moles of CH, and CzHa present at the various temperatures, a value for the over-all activation energy is obtained that is in close agreement with that calculated from the initial rate constants, k i . 4. Surface Efects. A reaction flask, identical in dimensions and volume with an unpacked bulb, was packed with 700 pieces of 3.0-mm Pyrex tubing 1.0-cm long. The surface-to-volume ratio of the unpacked bulb was 1.36/cm and that of the packed bulb was 16.7/cm. The packed flask was conditioned with a carbon coating similar to that of the unpacked flask before using. Reproducibility of results was excellent Volume 71,hTumber7 June 1967
C. E. WARINGAND REMOPELLIN
2048
in both flasks a t all temperatures. Table V summarizes the results of runs conducted in packed and unpacked vessels a t 401". The data indicate that the thermal decomposition of DMM is homogeneous in nature. 5. Products of Reaction. The reaction products analyzed were methane, ethane, propane, hydrogen, ethylene, and other unsaturated hydrocarbons. Repeated analyses, employing chemical, mass spectrographic, and vapor phase chromatographic techniques, showed that the only gaseous products present to any appreciable extent were methane and ethane. Higher hydrocarbons and olefins were found to be present only to the extent of 1-275. Table V"
s/ v
P(DMAfI),
1.36/cm 16.7/cm
mm
l/tiz.s%
l/tzs%
42.0 0.0520 0.0251 43.2 0.0528 0.0256
l/tw%
m/Pi
Pf/Pi
0.0092 0.0091
1.95 1.84
1.42 1.41
' l/tz% = reciprocal of the time for a given percentage increase over the initial pressure.
Table VI presents the gaseous products from the thermal decomposition of DMM at various time intervals. These time intervals were chosen so that the analyses were made at times corresponding to pressure increases of 25, 50, and 75% of the initial pressure. Analyses were also performed a t time intervals corresponding to the pressure maximum, p m , to n decrease in pressure to 70% above the initial pressure, and to the final pressure. By so doing, representative analyses were obtained over the curve shown in Figure 1. It is readily apparent from the table that the agreement between the chemical and mass spectrographic analyses is good. The data further indicate that, a t a Table VI: Mole Per Cent Gaseous Products from the Decomposition of 30 mm DMM a t 401.0' Time,
Yo Ap/pi
Be0
CH4
% CiHa
%
%
%
CaHe
CdHio
GHa
... ... ... ...
...
Chemical Analysis
56.4 43.2 ... 41.7 ... 58.1 57.8 42.0 ... 57.4 42.3 ... 58.4 41.3 ... 59.6 40.0 ... Mass Spectrograph 0.50 57.8 42.0 Trace 0.97 58.2 41.0 1.0 0.41 59.1 40.0 1.0
0.25 0.50 0.75 0.97 0.70 0.41
38 108 270 840 3,300 15,000 108 840 15,000
The J o u r d
of
Physical Chemdetry
...
...
... ...
... ...
...
...
...
2.0 1.5
2.0 1.8
given temperature, the CH4/C2Haratio is fairly constant over the entire pressure range. There is, however, a slight increase in [CH,] as the reaction proceeds. The decomposition of DRlAl also produced two solid products. Both were evident in the glass tubing leading out of the reaction furnace. One was finely dispersed mercury droplets that appeared during the phase of the decomposition when the pressure was decreasing from the maximum. The other was a yellow, translucent film which collected on the exit tube from the reaction flask outside the furnace when the gaseous products were being pumped out. A qualitative analysis of this material showed only carbon, hydrogen, and mercury to be present. It was not established as to whether the mercury was chemically bound or merely occluded. However, in view of the fact that the substance was highly colored, the assumption that the material was an unsaturated hydrocarbon polymer seems not unwarranted. A similar polymeric substance has been reported by Laurie and Long2 and Cattanach and Long2 but not found under the experimental conditions employed by Kallend and Purnell.' 6. Efect of Nitric Oxide. Previous have reported that nitric oxide has little or no effect on the rate of decomposition of DMM. Contrary to their observations, we find that nitric oxide produces several marked effects on this reaction. While nitric oxide does not change the general profile of the time-pressure curve in Figure 1, the pressure maxima, pm, and the final pressures, pr, were much higher in the fully inhibited reaction. For example, with 20 mm KO in the presence of 30 mm DMM a t 401.0", p,/pi = 2.48 and pJpi = 1.70, as compared to 1.96 and 1.41, respectively, for the uninhibited reaction a t the same temperature. For a given concentration of NO (30 mm) the p,/pi ratio decreases with increasing initial pressure of DMSS above 20 mm pi. This behavior argues that KO is not merely acting as an inhibitor but, in addition, is involved in some sort of chemical process. I n Figure 6 the ratio of the inhibited to the uninhibited rate constants is plotted as a function of the partial pressure of Y O for three different initial pressures of DMM a t 401.0'. It is seen that at all initial pressures of DMM, the limiting rate is attained a t NO = 15 mm and that the limiting value is unchanged out to approximately 85 mm partial pressure NO. Although the extent of the maximum inhibition is only lo%, this value is constant and reproducible a t all temperatures investigated. Since inhibition curves based upon rates of pressure change are not indicative as to the mechanistic effect of
THEKINETICSAND MECHANISM OF
THE
2049
THERMAL DECOMPOSITION OF DIMETHYLMERCURY
completely quenching the principal ethane-producing step. The course of the NO-inhibited reaction wa8 also followed by making careful analyses of all the products of reaction at various time intervals from the beginning to the end of the decomposition by both chemical and mass spectrographic methods of analysis. The results are presented in Table VIII.
6 1.1
2
s 0.9
0
16
40
60
80
100
Table VIII: Mole Per Cent of All Gaseous Products from the Decomposition of 30 mm DMM in the Presence of 30 mm NO a t 401.0"
PNO, mm.
Figure 6. Retardation of decomposition rate, I,&%, by nitric oxide a t 401.0" for: 0 , 15 mm; 0, 30 mm; 0,60 mm of dimethylmercury.
the inhibitor, the absolute amounts of methane and ethane were determined as a function of time for the NO-inhibited reaction. These data are presented in Table VII.
Table VII: Partial Pressures and Mole Per Cents of CH, and CgHa from the Decomposition of 30 mm DMM in the Presence of 30 mm NO a t 401.0" Time,
PCH4.
mm
sec
0 40
60 120 240 600 1,200 1,800 3,300 7,500 12,000
0 4.8 7.3 10.4 14.4 17.3 20.1 20.8 21.3 21.2 21.8
%Ha,
mm
% CHd
% CtHs
0 0 0 0.1 0.4 0.7 0.2 1.3 1.1 1.0 0.7
0 100 100 99.5 97.3 96.2 99.0 94.2 95.1 95.6 97.0
0 0 0 0.5 2.7 3.8 1.0 5.8 4.9 4.4 3.0
A comparison of these data with those in Table I shows that nitric oxide produces two rather startling effects upon the reaction mechanism. First, it is seen that, within the limits of experimental error, the concentration of methane at any given time interval is essentially the same in both the uninhibited and inhibited reactions. This clearly indicates that the reaction producing methane is unaffected by NO. Second, in the inhibited process, the ethane concentration is reduced to zero initially and only small amounts are generated in the latter stages of the decomposition. This argues strongly for the fact that NO is effective in
0.66 2.0 4.0 10 20 30 55 66.6 110 125 200 300
93.7 75.0 62.4 41.0 32.4 23.4 23.3 22.5 17.3 18.7 18.8 17.8
6.0 24.3 33.8 50.2 53.5 54.0
51.5 51.8 55.5 55.2 56.2 58.4
...
...
0.2 1.0 1.4 0.6 3.5 2.9 0.9 2.4 2.8 1.8 0.7
2.6 6.0 8.1 6.6 7.0 8.2 7.7 6.7 3.6 2.8
...
... *.. ...
... ... ...
1.0 4.1 9.3 11.6 13.1 10.8 11.5 12.2 13.6
0.2 0.4 2.8 3.6 3.6 6.0 5.3 6.5 6.8
Water was also identified as one of the reaction products, but since a quantitative analysis for water was not possible, it was not included in the table. A total of five mass spectrograms were made on the nitric oxide inhibited decompositions over five different time intervals from 2 to 110 min reaction time. In the early stages of the decomposition a fairly large peak was observed at mass number 45. At longer time intervals it rapidly diminished and was not observed after 20 min. From the nature of the reaction involved, it seems that the only compound that reasonably could correspond to a molecular weight of 45 would be fonnaldoxime, CH2NOH, or one of its tautomers. 7. E f e c t of Propylene. Propylene was found to have a pronounced inhibiting effect on the decomposition of DMM as measured by the rate of pressure increase. This is shown in Figure 7 where the ratios of the inhibited to the uninhibited rate constants are plotted as a function of the partial pressure of propylene. In contrast to the inhibiting effect of NO, it is seen that the extent of propylene inhibition is dependent upon the partial pressure of DMM. The pm/'pi and p J p i ratios for propylene were 1.86 and 1.28, respectively. Thus, these values differ Volume 71, Number 7 Juna 1067
C. E. WARINGAND REMOPELLIN
2050
&
1.1
s
culated on the basis of the gaseous products being 100% hydrocarbons.
5 $ 1.0
s.* 2 0.9 Y
Table X : Effect of Foreign Gases on the Decomposition of 30 mm of DMM a t 401'
._(
a 0
5
0.8
3 Y
.9 0.7 ._( 0
.u
-0
2 0.6 0
40
15
60
so
100
P C ~ H mm. ~ ,
Figure 7. Retardation of decomposition rate, I/~so%, by propylene for 15 ( O ) , 30 ( e ) and , 60 (0)mm of dimethylmercury a t 401.0'.
appreciably from both the NO inhibited and uninhibited reactions. Table IX presents the effect of propylene on the reaction products from the decomposition. These data indicate that the main effect of propylene is to reduce the concentration of ethane below that found in the uninhibited decomposition (Table I). They also suggest that the concentration of methane has been increased. Since it is known that propylene reacts with methyl radicals to form methane and allyl radicals, the true effect of propylene on the reaction's methaneproducing step cannot be evaluated. The data in Table IX also emphasize the fact that propylene is a much less efficient inhibitor of the ethane-producing step than NO.
Table IX: Mole Per Cents of CHI and C,He from the Decomposition of 30 mm DMM in the Presence of 30 mm Propylene a t 401.0' Time, sec
60 180 300 600 1200
-Chemical % CHI
73.9 74.7 74.9 76.5 ,
..
analysis% CzHa
26.1 25.3 25.1 23.5
..,
~ V p c - % CHd % CzHs
73.5 74.6 74.4 75.2 76.9
26.5 25.4 25.6 24.8 23.1
Mass -spectroscopy% CH4 % CzHs
... .. .
*..
77.8
22.2
... ...
...
...
...
8. Effect of Other Foreign Gases. I n addition to nitric oxide and propylene, the effects of hydrogen, nitrogen, and helium on the decomposition of DMM were also studied. Table X compares the effect of the foreign gases on the CH4-CzHe production, cal-
The Journal of P h y s h l Chemistry
Foreign gas, 30 mm
% CHI
% CzHa
CHdCzHs
None NO (at p d CZH6 (at pm) Hz Nz He
59.4 98.8 75.2 86.8 67.6 77.7
40.1 1.2 24.8 13.2 32.4 22.3
1.49 82.2 3.04 6.57 2.08 3.48
These data show that while propylene, hydrogen, helium, and nitrogen cause a decrease in the amount of ethane formed, they do not reduce its concentration in the dramatic fashion of nitric oxide.
Discussion The experimental evidence presented here indicates that DMRl thermally decomposes to give essentially only methane and ethane. These results are in agreement with those of a number of investigators2 who have studied this reaction both thermally and photochemically. The only other products found in the uninhibited decomposition were 1-2% of butane and butene. Attempts to find all of the products reported by Kallend and Purnell' were, under our experimental conditions, unsuccessful by three independent analytical techniques. Although a number of mechanisms have been proposed for the thermal decomposition of DMM, none successfully account for all of the experimental kinetic and analytical data. Taylor and Cunningham,2 for example, postulate that the ethane formed in this reaction is the result of (5)
2CH3 +C2Hs or possibly to CH,
+ CH,HgCH,
+C2Ho
+ Hg + CH3
(6)
Reaction 5 cannot be defended on two important counts. First, it can be seen in Table I that the concentration of ethane in the initial stages of the decomposition is high, being nearly equal to that of methane. Since the concentration of DhfM in the early stages of the reaction would be extremely high in comparison to the relatively small steady-state concentration of methyl radicals, it is difficult to understand how reaction 5 alone could account for the observed amount of
THEKINETICSAND MECHANISM OF
THE
THERMAL DECOMPOSITION OF DIMETHYLMERCURY
ethane, even though the activation energy for this process is zero or nearly so. The formation of ethane by reaction 6 avoids this criticism. Second, neither reactions 5 nor 6 can account for the observed inhibition effect of nitric oxide. Table VI1 shows that NO reduces the ethane concentration virtually to zero while the methane concentration is essentially the same as in the uninhibited decomposition. It is obvious, therefore, that NO is not removing methyl radicals and that most of the ethane formed, at least, must be produced by a mechanism other than reactions 5 or 6. Kallend and Purnell,‘ on the other hand, have proposed that ethane is formed according to CH3
+ CHzHgCH3
4
CH3CHzHgCHa
+ Hg + CH3 CzH, + CH3HgCHs +CzH6 + CHzHgCH3 CH3CHZHgCH3 +CzHs
(7)
(8) (9)
While this mechanism can successfully account for the effect of NO on the formation of ethane, it also fails t o account adequately for the observed quantity of ethane found initially. Certainly in the early stages of the reaction the chance of a CHZHgCH3 radical colliding with a methyl radical could be no greater than that of a methyl-methyl collision. Like the methyl radical, the chance of CHeHgCH3 initially colliding with anything other than a DMM molecule would be negligible. Further, if ethyl radicals are indeed formed as indicated by reactions 7-9, one would expect t o find detectable quantities of methyl ethyl mercury, C3Hs, and C4H10 in the reaction products. The authors report only the presence of C3H.3. It is evident, therefore, that any mechanism proposed for the thermal decomposition of DMM must be able t o account for the following salient experimental facts: (a) the rate of decomposition of DMM is of the first order, (b) the gaseous reaction products are essentially only methane and ethane, (c) the amounts of methane and ethane produced initially in the uninhibited decomposition are approximately equal, (d) the rates of formation of methane and ethane also follow first-order kinetics, and (e) in the nitric oxide inhibited reaction, the concentration of ethane is reduced virtually to zero while the concentration of methane remains essentially the same as in the uninhibited decomposition. A mechanism which successfully satisfies these criteria is postulated as CH3HgCH3 CH3
+ CH3HgCH3
5CH3 + HgCH3 CHd
+ CHzHg + CH,
(10) (11)
HgCH3
+ CH3HgCH3-% CzH6 2CH3
+ 2Hg + CH3
2051
(12)
+ ?tf 2CzH6 + R/!
(13) There is general agreement among investigators as to step 1. Step 2 is undoubtedly the only methaneforming reaction since the probability of a methyl radical colliding with any other particle than DMM in the initial stages of the decomposition is extremely remote. It also accounts for the observed presence of the yellow translucent solid, probably (CH2Hg),. Step 3 is postulated as being the principal ethaneproducing reaction. The postulation of a three-body process for the formation of ethane in step 4 is not unreasonable, especially at lower pressures. Since the data in Table V indicate the absence of any surface effects, the assumption is made that the third body, M, is a DMM molecule. In the event that objection might be raised against step 4 a t higher pressures, it will be seen that this step may be written as two successive bimolecular processes which give the same over-all third-order kinetics. Among the arguments strongly supporting this mechanism are the kinetic and energy relationships. If one makes the usual steady-state assumptions it can be shown that d [CH,]/dt = kl [DlKiII]
+
k3[HgCH3][DMRII] - 2k4[CH3I2[M] = 0 (14) and d[HgCH3]/dt = kl[DMRI] k3[HgCH3][DRfRf] = 0 (15) Adding eq 14 and 15 and remembering that [MI = [DMM] [C&] = (ki/k4)”*
(16)
[ H ~ C HI B= k i / h
(17)
and The rate of disappearance of DMM is then found to be -d[DMM]/dt = [2k1
+ IG~(ki/k4)”*][DR/IM]= dp/dt
(18)
Thus, the rate of decomposition of [DMM] is seen to be first order in agreement with that observed experimentally. I n like manner, the rates of formation of methane and ethane can also be shown to be first order, again in agreement with experiment. Volume 71,Number 7 June 1067
C. E. WARINGAND REMOPELLIN
2052
kcHk
= kz[CH3] [DRIIM] = k2(kl/k4)’/’[DMM]
(19)
and
+
l C C z ~ s= k3 [HgCHa][DMbLI.I k4
[CH3]2[MI = 2k1 [DMRII] (20)
It can further be demonstrated that the rates of formation of methane, kCHk1 and ethane, k + z ~ sCan , be equated to expressions 3 and 4, respectively k+H,
=
k 2 ( / ~ 1 / k * )= ~/~ 6.84
x
109e-32,200/RT sec-’
(21)
1012e-45,100/RT sec-’
(22)
and
kCzHa= 2kl
= 1.93
x
from which /cl = 0.97
x
1012e-459100’RT see-1
(23)
Since the strength of the G-Hg bond in DMM is given as 51.0 kcal/mole16 an activation energy somewhat less than this for is not unreasonable. From eq 21 and 23 and the fact that the activation energy for step 4 is zero, the activation energy for k2 is calculated to be 9.6 kcal/mole. This value is in excellent agreement with other hydrogen abstraction reactions by methyl radicals. A further check on the validity of the experimental data can be made by calculating the preexponential factor fork,. The activation energy for the reaction
+
for the unimolecular collision frequency. For the reaction to be third order it is necessary, of course, for k b >> k,[M]. Since [MI is less than mole/l., this is indeed the case. The mechanism given in steps 10-13 does suggest that slightly more ethane than methane should be formed. The fact that somewhat less ethane is actually found is explainable on the basis that some of the HgCH3 formed in step 1 decomposes into CH3 and Hg before it can react with DMM. The contribution of step 4 to the total ethane concentration is very small, though measurable, as will be shown later. Another argument in favor of this mechanism is that it can adequately explain the inhibition effect of NO. As previously mentioned, in view of the essentially unchanged methane concentration, NO cannot be removing methyl radicals directly. If, instead, NO removes the large radical, HgCH3, then step 3 would be eliminated and the only ethane formed would be by step 4. Table VI1 indicates that a small amount of ethane is, indeed, formed in the latter stages of the decomposition. The net effect of the elimination of step 3 is to give a steady-state concentration of methyl radicals of [CH,] = (k1/2k4)1’2
The rate of formation of methane in the inhibited reaction then becomes
+
CH3 CHaOCH3 +CH4 CHzOCHs (24) is given as 9.5 kcal/mole7 and 3.2 X lo8 I./mole sec for the frequency factor. From eq 21 0.97 x 1012e-45t’00/RT(3.2 x lOBe-Q5W/RT )2 k4 = - (6.84 X 109e-32*200/RT2 ) 2*1 x log1-2/mo1e2
(25)
This value also agrees well with that of 1.1 x 109 1.2/ mole2 sec givens for the collision frequency of a threebody process. It has been pointed outg that the mechanism for methyl radical combination in step 4 may be written as two successive bimolecular processes. ks
2CH3 E C2&*
(26)
kb
C2H6*
+ M k., CZH6 + M
(27)
k, is givens as 2.2 X 1Olo l./mole sec and k, is of the order Of the bimolecular collision frequency, or about 2 X 10” I./mole sec. Thus, kb can be calculated from k4 = k,k,/kb = 2.1
x
109
1 . ~ / ~ ~ ~(28)~ z
to be about 2 X 10l2sec-’, which is a reasonable value The Journal of Physical C h a k t r y
(29)
kcHk
= k2(ki/2k4)1/2[Dl\IM]
(30)
If one
eq 30 with eq 19 it is immediately apparent that the rate of formation of methane should be virtually the same in the uninhibited and inhibited reactions. This conclusion is supported by the data in Tables I and VII. It can be further argued that if NO reacted with methyl radicals to any appreciable extenti it would be expectedthat the equilibrium CHI
+ NO
CH3NO
(31)
would be shifted strongly to the left a t these temperatures. Under these conditions, the steady-state concentration of methyl radicals and the concentration of NO would be essentially unchanged. On the other hand, if NO reacts according to HgCH3
+ NO +NOHgCH3
(32)
(6) T. L. Cottrell, “The Strength of Chemical Bonds,” Academic London*1958i 210. (7) 8. W. Bensen, “The Foundations of Chemical Kinetics,” McGraw-Hill Book Co., Inc., New York, N . Y., 1960, p 391. (8) See ref 7, p 157. (9) w e are grateful to a referee for this suggestion.
THE KINETICSAND MECHANISM OF
THE
NOHgCHa +CHzNOH
2HCN
THERMAL DECOMPOSITION OF DIMETWLMERCURY
+ Hg
+ HzO CO + NZ + polymer
(33)
CHzNOH +HCN
(34)
+ 2H20
(35)
-3
then NO should disappear as the reaction proceeds. Table VI11 shows that this is the case experimentally. It is also of interest to note that the pressure maximum in the NO-inhibited reaction corresponds to the maximum concentration of HCN. Additional support for the fact that NO reacts preferentially with large radi-
a
25
50
15
Time, min.
100
cals rather than methyls is found in the studies of Hobbs'O and Smith and Hinshelwoodll on the pyrolysis of diethyl ether. One final argument can be offered in support of the mechanism proposed in reactions 10-13. It was postulated here that most of the ethane was produced by reaction 12 rather than 13. For this assumption to be valid, after the initiating step DMM should disappear primarily through two independent mechanisms : one involving CHI, the other involving HgCH3. If NO removes the radical responsible for producing ethane, namely HgCHa, but does not interfere with the methane-producing step by attacking methyl radicals, then the disappearance of DMM is due primarily only to reaction 11. Since the rate of this reaction has been shown by eq 19 and 30 to be the same in both the uninhibited and inhibited processes, one would then expect the rate of disappearance of DMRI t o be slower in the inhibited reaction. Figure 8 shows this t o be the case.
Acknowledgment. The authors gratefully acknowledge the support of this work in part by the U. S. Naval Ordnance Test Station, China Lake, Calif., under Contract No. N123~-60530~(5150a).They also wish to thank Drs. S. R. Smith and C. W. David of this department for their help in the interpretation of the mass spectrographic analyses and for their valuable suggestions regarding mechanisms. ~
Figure 8. Rate of disappearance of 30 mm of DMM a t 401': uninhibited, 0;inhibited (30mm NO), 0.
2053
~~
~
(10) J. E.Hobbs, Proc. Roy. SOC.(London), A167, 456 (1938). (11) J. R. E. Smith and C. N. Hinshelwood, ibid., A180,468 (1942).
V o l m 71, Number 7 June 1067
