Shock Tube
The Journal ol‘ Physical Chemistry, Vol. 83, No. 2, 1979 297
Studies of Monomethylamine
Initial Shock Tube Studies of Monomethylamine Ernest A. Dorko,” Nicholas R. Pchelkln, John C. Wert, 111, and Gwhard W. Mueller Department of Aeronautics and Astronautics, Air Force Institute of Technology, Wright-Patterson Air Force Base, Ohio 45433 (Received July 24, 1978; Revised Manuscript Received November 17, 1978) Publication costs assisted by the Air Force Institute of Technology
The decomposition of dilute mixtures of monomethylamine (MMA) in argon were studied in the reflected wave of a shock tube, Kinetic measurements of both the decomposition of MMA and the formation of ammonia, one of the postulated reaction products, were made by use of infrared (IR) emission techniques. The decomposition was studied over the temperature range 1275-2400 K and the total pressure range 1-10 atm. Both MMA decomposition and ammonia formation were found to obey the Hinshelwood-Lindemann model for a unimolecular reaction. Arrhenius parameters were obtained over the experimental temperature and pressure ranges and the extrapolated parameters for high and low pressure llimits were obtained. In the high pressure limit E, = 48.15 kcal/mol and log A = 10.84. A t low pressure E, = 35.14 kcal/mol and log A = 13.50. Various reasonable mechanisms are discussed for the decomposition and evidence is offered to indicate that the predominant mechanism is C-N bond scissure followed by a short free-radical reaction sequence.
Introductioin There is current concern about the high level of NO, pollutants which form during the combustion of some coals and shale oils which contain organo-nitrogen compounds. It seems that the reactions which lead to the formation of NO, pollutants from these organo-nitrogen compounds occur much more readily than do the reactions which lead to the fixation of atmospheric nitrogen by means of the Zeldovich mechanisms.’ This knowledge has prompted the study of the inechanistic routes by which such transformationis occur. The pyrolysis of dilute mixtures of monomethylamine in argon was studied as the first step toward modeling these routes. Monomethylamine (MMA) was selected as a model compound because it is in many ways the simplest, of organo-nitrogen compounds, it is readily available, and it is obviously analogous to ethane, a hydrocarbon whose combustion has been studied previously. Even with this relatively simple compound, however, three distinct decomposition reaction mechanisms can be reasonably postulated. In initial studies by Hidaka, Saito, and Yamamura,2they postulated sciwure of the CN bond to form the methyl and amide radicals followed by a radical chain reaction. The mechanism is postulated as follows: CH3NHz CH3 -I- CH3NHz
-kl
+ NH2 CzHG + NH2 CH4 + CH2NHz
CH3
(k2)
-+
(3 1
NIIZ
+ CH3NH2
CH3
+
--
+ NZH4
(3)
+
kl
CHBNH2 --CH, NH2 CH3 + CH3NH2
kz
CH4 + CHzNHZ
+
k5
CHzNH2 -* CHZ=NH
H
k6
+ CH3NH2 H
+H
Hz + CHzNHz
-3
+ CH,
k7
CH4
(1) (2)
(5)
(6)
(7)
By use of the steady state approximation the rate expression for the disappearance of MMA is
which is a combination of first- and second-order reactions with respect to MMA concentration. The final CN bond scissure reaction may be postulated as follows: CI-13NHZ
(2) k3
tube study in conjunction with vapor phase chromatography3details the decomposition of ethane. The data of this report correlate well with the assumption of unimolecular reaction kinetics for the initial step. A similar reaction sequence for MMA would proceed as follows:
kl
CH,
+ CHSNHZ
CH,
+ CHzNHz
CH3 + NHZ
k2
CH4
+ CHzNH2
(2)
CH4
+ CHz=NH
(8)
--+
k8 ---+
(1)
k4
CzH6 (4) 2CH, Based upon this mechanism a simplified rate expression for the disappearance of MMA was developed as (-d/’dt)[CH,NHJ
NH2 + CHBNH2 NH2 .f CHzNHz
2kz(2kl/k4)1’2[CH3NH213/2 (I)
This expression indicates a 312 order power dependence of the reaction on monomethylamine concentration. The CN bond scissure may also be considered analogous to the unimolecular bond scissure of ethane followed by a radical chain mechanism. A recent single-pulse shock 0022-3654/79/2083-0297$0 1.OO/O
CI32=”
k9
NH,
-t
CHzNHz
(9)
NH3
+ CH2=NH
(LO)
+
kl,
kll +
Hz
+ HCN
Also by use of the steady state approximation, the rate expression for the disappearance of MMA is (-d/dt)[CH,NH,] 0 1979 American Chemical Society
= 2k1[CH3NHz]
(111)
298
The Journal of Physical Chemistry, Vol. 83,No. 2, 1979
which is a first-order reaction in MMA. The second reaction mechanism which can reasonably be postulated involves the formation of a long-lived intermediate from which HCN is then formed: CH3NHz
-+
+ Hz
CH2=NH
CHZ=NH
+
HCN
+ Hz
(12) (13)
This mechanism was proposed recently by Smith and Sawyer4to explain the decomposition of MMA mixed with helium in a flow reactor. They found methane, ammonia, nitrogen, traces of hydrogen cyanide, and solid hexamethylenetetramine (l),which they postulate is produced
1
from the imine intermediate. A t higher temperatures the imine intermediate breaks down to form HCN as shown in eq 13. The experimentally determined rate expression is (-d/dt) [ CH,NHz] = k[ CH3NH2]0.28
(IV)
The last mechanism to be postulated is as follows: CH,NHZ CHZ + NH3 (14) +
2CHz
-+
CHZ=CHZ
(15)
There is some evidence to show that this mechanism can reasonably be expected to occur. A recent shock tube study by Meyer and Wagner5 on hydrazine presents evidence for hydrogen transfer as follows: NzH4 NH3 + NH (16) -+
At this point it should be noted that the first step in the three different postulated reactions is a unimolecular decomposition. However, the reaction pathway is significantly different in each case. The three pathways are shown in eq 17. Pathway a would be followed if C-N bond
scissure occurred; pathway b if a four-center intermediate were involved; and, pathway c if a three-center intermediate were involved. After the first step, there are substantial differences in the subsequent reaction steps and in the products which are formed. The present report describes the results of tests performed in a shock tube on very dilute mixtures of MMA in argon. These tests were performed in an effort to determine the dependence on temperature and pressure of the rate of the first step and to attempt to decide, if possible, which of the pathways postulated in eq 17 is the dominant one. Analysis of the data was performed with the assumption that the basic Hinshelwood-Lindemann (HL) mechanism6 is an adequate model. In terms of the HL model
Dorko et al.
where k,, kd, and k, are the rate constants for activation, deactivation, and reaction, respectively, and [MMA] and [MI are MMA and total gas concentrations, respectively. Equation V thus defines the experimentally determined rate constant, huni, in terms of the HL model. Experimental Section The shock tube used in these experiments was fabricated from a piece of 3411. i.d. stainless steel (SS304) tubing with 3/s-in. walls. The tubing was shaped on a mandrel to produce a cross section at the reaction end which had two parallel, flat sides with a width of about 1 in. The flat sides made it convenient to attach instrumentation and the absence of corners minimized boundary layer effects. The tube was 23 ft long with a 16-ft driver and a 7-ft driven section. The tube has been described more fully elsewhere.’ MMA test gas mixture (0.25, 0.5, 1, and 2%) were prepared from MMA (98.0%) and argon (99.995%), both purchased from Matheson Gas Products. Shock parameters were calculated for the specific gas mixtures from the initial shock velocity assuming frozen chemistry. Heat capacities for MMA were obtained from the data contained in NASA program 273 and from thermodynamic tables.8 The total gag concentration behind the reflected shock ranged from 14.00 to 114.7 pmol/cm3. The temperature range was 1275-2400 K and the total pressure ranged from 1 to 10 atm. A total of 400 shocks were used in the analysis. Kinetic measurements were obtained for the decomposition occurring behind the reflected shock by detection of IR emission through a CaFz window placed in the flat wall of the shock tube 1 2 mm from the end flange of the driven section. This window was fabricated by the Industrial Lens Co., Dayton, Ohio, and mounted in a stainless steel adapter. The window surfaces were made flat to one-half wavelength and parallel to 10” of arc. Infrared emission was monitored by an optical system which consisted of the focusing lens, pinhole/filter assembly, and indium antimonide (InSb) detector. The lens was made from CaFz and had a focal length of 25 cm. The InSb detector was a Barnes Engineering Type J-10 (liquid nitrogen cooled) used in conjunction with a Perry Associates Preamplifier Model 720. The amplifier signal was fed into a Tektronix Type 551 oscilloscope which was triggered by an output signal from the last velocity station. A signal originating from a Tektronix Type 180A time mark generator was fed into the lower beam of the oscilloscope. A photographic record of the oscilloscope trace and the time marks was made in each case with Polaroid film. The oscilloscope traces of the IR emission intensity vs. time were converted to numerical data with a traveling microscope. Series of experiments were run in which a 3.375-pm filter (fwhm of 0.20 pm) from Optical Coatings Laboratory was inserted in the pinhole/filter assembly of the InSb detector system in order to monitor the decomposition of MMA. The emission at 3.375 pm is due to the vz band of MMA.’ Numerical data were obtained from each oscillogram by measuring the relative intensity vs. time by use of the following procedure. First, a reference baseline, which is an extension of the portion of the trace just prior to the arrival of the reflected shock, was drawn on the oscillogram. Next, a smooth curve was drawn through the trace obtained after passage of the reflected shock wave. The initial time was chosen as the time mark coincident with
Shock Tube Studles of Monomemylamine
The Journal Of physical Ctmmkby. Vol. 83. Na. 2, 1979 299
I
Figure 1. Typml oscillogram fa IJe decomposaion of MMA. T = 1678 K. p = 4.58 alm. [MI = 33.2 yrnol/cm'. l%MMA in argon. writing speed from left 10 right = 50 &cm. time marks every 10 11s. vertical scale 5 mV/cm.
Figure 3. Typical oscillogram f a the emission at 5.230 ym. T = 1617 K. p = 4.33 atm. [MI = 32.6 11moI/cm~,1% MMA in argon. writing speai from left to right = 100 @s/crn.lime marks e v w IO ys, vwlical scale 2 mV/cm.
different method of data reduction was used. The baseline was drawn through the zero intensity portion of the oscillogram and a smooth curve was drawn through the remainder of the trace. This curve was then extended toward infinity and a line tangent to the curve at this "infinite" point was constructed back to time equal to zero. This line represented the maximum intensity (I-) for each oscillogram. The relative intensity (ImI) was then calculated by use of the following formula:
(VI) is the trace intensity at any time between 0 and Le, = (Imaa -
Figure 2. Typical oscillogram for the formation of NH,. T = 1635 K. p = 4.41 alm. [MI = 32.9 ymo1/cm3. 1 % MMA in argon. writing speed from left lo right = 100 yslcrn. lime marks every 10 ys, vertiml scale I
mv/cm.
the maximum intensity of the oscillogram. See Figure 1 for a typical MMA oscillogram depicting the baseline, smooth curve, and time marks. A Gaertner Model 1150 traveling microscope modified to meet present requirements was used to measure the relative intensity vs. time of the oscillogram. The hand drive on the microscope was mechanically coupled through a 4:l gear reduction to a ten-turn linear potentiometer, which was electrically connected to a digital voltmeter through electronic circuitry. This circuitry permitted the setting of an arbitrary zero point and then relative readings of the position of the microscope with respect to the arbitrary baseline could be obtained. These numerical results were recorded as relative intensity, which is proportional to MMA concentration vs. time. A plot of the logarithm of the relative intensity" vs. time yielded a straight line, the negative slope of which was defined to be kunc Other series of experiments were run in which a 2.886-pm filter (fwhm of 0.22 pm) from Optical Coatings Laboratory was inserted in the pinhole/filter assembly of the InSb detector system to monitor the formation of ammonia, NH3. The emission at 2.886 pm is due to A typical oscillogram is shown in Figure 2. Inasmuch as the emission intensity is increasing with time, a slightly
Lned/Irnax
where I,. infinity. In addition to the series of tests described, another series was run with a 5.230-pm filter (fwhm of 0.36 pm) from Optical Coatings Laboratory inserted in the pinhole/filter assembly of the InSb detector system. A typical oscillogram is shown in Figure 3. Because of the complexity of the process which gave rise to the oscillogram, no simple data reduction and analysis were possible.
Results and Discussion If the HL mechanism holds, then an exponential decay of the reactant concentration is to be expected. The emission at 3.375 pm due to the vz band of MMAg showed this expected exponential decay, which is demonstrated by the linearity of a plot of the logarithm of the relative intensity vs. time. The negative slope of the line thus obtained is kUni. Arrhenius curves were prepared from the values of kuni. The analysis was performed in the following way. A series of shots was made in which the total gas concentration was kept constant and the MMA concentration was varied by as much as a factor of 8. A least-squares line through the data points was made by a computer program READN,'* which was developed for this study and the results are depicted in Figure 4 as an Arrhenius plot. This leastsquares line has a standard deviation of 0.089 and the correlation factor for the data is 4.98. It is concluded that there is no dependence of the rate constant, kUni,on the concentration of MMA. An error analysis of each memher of the series made independently of the other series members was also prepared by least-squares fit of the data in the Arrhenius plots. The conclusion from the latter analysis is that all the curves were within 1 u of the total data curve over the temperature range of this study and there is no power dependence of kVnion MMA concen-
300
The Journal of Physical Chemistry, Vol. 83, No. 2, 1979
Dorko et al.
TABLE I: Arrhenius Parameters for Monomethylamine Decomposition
c)
1. HETHYL RnINE I N R R * m 1. H E T H I L R n l Y f 1Y RRGM . 5 1. HETHYL M I N E IN RRGMi . 2 5 1. HETHIL MINE I N RRDDN
2. I.
0 A
+
log A , s-' 15.32 27.50 53.04 83.73 114.7 low press. high press. a
40.79 i 41.17 i 44.88 i 47.07 i 47.39 f 35.14 48.15
0.61 2.25 1.62 1.46 0.71
9.27 i 9.48 i 10.20 li: 10.62 ? 10.70 f 13.50a 10.84
0.14 0.62 0.37 0.33 0.16
Units of A are cm3/(mol s).
n K c)
4.80
5.20
6.00
5.60
5.10
6.80
10.~4 / T
7.20
7.60
8.00
8.10
Figure 4. Arrhenius plot of kUni (MMA decomposition) for a constant total gas concentration of 32 f 3 pmol/cm3.
'
/
"1
4 .IHR O E C C H P O S I T I O N
T
L B O
I
-
5 2 0
!
1.60
6.00
----
6.10
6 E O
7.20
10.~4 / f
Figure 5. Arrhenius plots of k,,, (MMA decomposition) at various constant total gas concentrations.
tration. Furthermore, it can be concluded that the decomposition is first order in MMA concentration. In the next series of tests a determination of the dependence of hunion total gas concentration was made. In the low pressure regime huniis defined as (VIII) Then at two different total gas concentrations the power dependence, p, of hunion [MI can be determined. The following analysis is used: The concentration [MI, is proportional to [MI, by a constant factor a , i.e. [MI, = a w l 1 (IX) kuni,
= ha[MIIPaP
(XI
Then the ratio of unimolecular rate constants is huniz/kuni, = a'
(XI)
Taking the natural log of eq XI and solving the equation for ,B yields (XIII p = (In kuniz- In kuni,)/ln a According to the Hinshelwood-Lindemann mechanism,
0 is equal to unity in the low pressure regime and to zero
2
3
4
5 6 78$
IO'
2
3
Figure 6. Plot of log k,,, vs. log Pat various constant temperatures.
in the high pressure regime. The Arrhenius plots for the decomposition at five different total gas concentrations are shown in Figure 5 . The curves were obtained by a least-squares analysis of the data.12 It can be noticed that there is somewhat of a fanning out of the curves. This effect is expected if the system obeys the HL theory. The reason for this is that as the pressure is increased the activation energy of the reaction (and therefore the slope of the Arrhenius plot) increases until the high pressure limit is reached. Notice that this fan effect is most pronounced at low pressure and is hardly noticeable as the system reaches its high pressure limit. In addition to the change in slope, the separation between curves decreases as the pressure is increased. This result is as expected with an HL reaction. At a temperature of 1667 K the log difference between the two lowest pressure curves is 0.197 while the log difference between the two highest pressure curves is 0.044. These differences correspond to total concentration power dependence, p, of 0.78 and 0.32 at the low and high pressure experimental extremes, respectively. The values of the Arrhenius parameters for MMA decomposition obtained
The Journal of Physlcal Chemistry, Vol. 83,No. 2, 1979 301
Shock Tube Studies of Monomethylamlne
TABLE 11: Arrhenius Parameters for Ammonia Formation
.--
[MI
9
m~oI/cm3
- .___
27.50 53.04 83.73
E,, kcd/mol 40.75 44.56 46.48
* f
*
1.49 2.09 1.75
log A , s-' 9.28 9.87 10.37
* 0.34 f 0.47 * 0.39
for these curves are given in Table I. Figure 6 displays plots of log h,,, vs. log P for various constant temperatures. The data for these plots were obtained from the Arrhenius parameters in Table I. At the high pressure limit in the HL mechanism the theoretical variation of k, with concentration or (from the ideal gas law) with pressure is equal to zero. It is evident that the experimental variation of h,, with concentration is approaching this zero limit. At the low pressure limit the HI., mechanism predicts a linear dependence of h,, on concentration. It can be seen from Figure 6 that as the temperature iincreases the slopes of the lines approach, but do not reach, unity. This indicates that the experimental observations were made in the intermediate to high pressure regime. The low pressure regime was beyond conditions thiat could be attained in the present experiment. By use of the data points from this figure, the values of the Arrhenius parameters for the high pressure regime and for the lowest experimental pressure regime were calculated. The values for the Arrhenius parameters for the high and low pressue regimes are shown in Table I. The decrease in E, from high to low pressure (13.01 kcal/mol) corresponds to a decrease of about 4RT in the temperature range under consideration. This decrease is not as expected from the simple HL theory. According to the theory 1he difference between high and low pressure activation energies is6 E -- Eo = ( S - l)RT (XPII)
I 111
60
1 10
3160
51 10 ICn.4
7
8'60
7
7
70
-
-
3 60
-
7
P 03
8 10
/ T
Figure 7. Arrhenius, plot of k,,,' (NH, formation) at various total gas concentrations.
3
%
a
where 5' is a number usually taken to be 2/3 the number of vibrational fundamentals, which for MMA is (2/3)(15) or 10. Thus, an activation energy difference of about 27 kcal is expected between the high and low pressure rer 53.04 gimes. The conclusion from this discussion is that the low pressure regime has not been reached under the experimental conditilons used. For a molecule such as MMA with seven aloms this is not an unexpected result. The low 20 60 $0 20 1 40 10n.4 / 1 pressure regime for such polyatomic molecules may not occur before the pressure is well below a t m o ~ p h e r i c . ~ J ~ Figure 8. Comparison of the Arrhenius curves for MMA decomposition and NH3 formation a l [MI = 53.04 pmol/cm3. The log difference From these results it can be concluded that the debetween the curves corresponds to a ratio of rate constants of 2.13. composition is unimolecular and that the HL mechanism is at least an adequate model to use for its analysis. In From eq XIV it is obvious that y = hunl/huni'.Figure 8 order to make further inferences regarding the dominant shows the Arrhenius plots for MMA decomposition and mechanistic pathway the following results need to be for NH, formation under identical experimental conditions considered. for a constant total gas concentration of 53.04 pmol/cm3. Plots of the logarithm of Ire,for NH, (eq VI) vs. time The value of y is 2,13. Additional comparisons at total gas produced straight lines. These results indicate that NH, concentrations of 27.05 and 83.78 pmol/cm3 gave values formation follows unimolecular kinetics as does the defor y of 1.58 and 1.78, respectively. The conclusion from composition of' MMA. Rate constants were determined this data is that MMA decomposes at approximately twice from the slopes of the lines and were designated h,,'. The the rate at which NH3 is formed. least-squares Amhenius plots obtained for the formation of NH3 at different total gas concentrations are shown in General Discussion Figure 7, A fanning out of the curves similar to that The experimental results need to be discussed in the encountered with the MMA decomposition is apparent. light of the reasonable mechanisms presented earlier. This The values of the Arrhenius parameters are located in discussion is an attempt to decide if one of the mechanisms Table 11. is dominant and if' so, which one. Given the okiservation that NH3 formation is unimoReferring to eq 17, the following conclusions can be lecular it can be shown14 that the rate equation for its drawn from the evidence presented earlier. The experiformation is the following: mentally determined rate of formation of NH, as one-half d[",I/dt = k"n,'[MMA] = (h,n,/r)[MMA] (XIV) the rate of MMA decomposition precludes the hydrogen
-1
:
% 1 7 4 10
1 BO
5
b
I
6 00
6
6 80
7
60
7 -
L1 00
8
302
The Journal of Physical Chemistry, Vol. 83, No. 2, 1979
TABLE 111: Heats of Formation for Selected C o m p o u n d s A
%,a
substance
kcal/mol
substance
CH, CH, CH,CH, CH,CH, CHCH
-17.889 34.0 - 20.236 12.496 54.194 -11.4
NH, CH,CH, CH,NH CH,NH HCN
"3
A&,a kcal/mol
47.2 -6.7 45.4 9.02b 31.2
E s t i m a t e d from t h e a Reference 1 6 except as n o t e d . heats of f o r m a t i o n of CH,NH, a n d HCN. See r e f 14.
transfer or three-center intermediate of pathway c (eq 14) since that mechanism would require a rate of formation of NH3 equal to the rate of MMA decomposition. The four-center mechanism of pathway b (eq 12-13) involves the formation of a long-lived imine and would give as products only hydrogen and HCN. Since NH, is formed in the present experiment as evidenced by the emission at 2.886 pm9 and since preliminary evidence from gas chromatographic analysis indicates the presence of substantial quantities of methane, this mechanism cannot be the dominant one for the decomposition of MMA. Having eliminated the three-center and four-center reaction pathway, the only remaining reasonable pathway is the one involving C-N bond scissure as the first step. Let us then consider the three free-radical mechanisms postulated earlier. The decomposition is first order in MMA concentration and the power dependence on total gas concentration varies from 0.32 to 0.78. These two results are strong evidence to eliminate two of the three postulated free-radical mechanisms. The model proposed by Hidaka, Saito, and Yamamura2 which leads to a 3/ 2-order rate expression is not supported by the established first-order dependence, although the values of E, and A are comparable to present results. In addition, the mechanism does not show NH3 as a product. The firstorder dependence also does not support the ethane analogous reaction which is a combination of first and second orders. On the other hand, the knowledge that y = 2 argues for the dominance of the t,hird free-radical mechanism. From eq 1, 2, 8, 9, 10, and 11 it can be seen (and it can be mathematically established) that under the steady state assumption,14 MMA decomposition must occur at twice the rate of NH3 formation. The fact that y 5 2 is strong evidence in favor of the third radical chain mechanism. Additional evidence to indicate the dominance of this mechanism follows. The following heats of reaction were calculated for some of the postulated reaction steps using the values shown in Table 111. The three reaction steps to be considered are given in eq 5, 8, and 10. Their heats of reaction are calculated to be +15.7, -88.3, and -94.9 kcal/mol, re-
Dorko et al,
spectively. Since reactions 8 and 10 are highly exothermic and reaction 5 is endothermic, reactions 8 and 10 would be strongly favored over reaction 5. This result would tend to eliminate the ethane analogous decomposition of MMA from consideration. At this point it should be noted that CH2NH2 and CH3NH have been used interchangeably. This is due to the fact that the activation energies for their formations are 8.7 and 5.7 kcal/mol, re~pective1y.l~ Therefore, there would be no noticeable preference for the formation of one over the other. Another point in favor of the interchangeability is that the bond strengths of C-H and N-H in CH3NH2are 972 and 9216 kcal/mol, respectively. A final consideration is pertinent. The high pressure E, (48.15 kcal/mol) is substantially lower than the estimated 79 kcal/mol C-N bond strength in MMA.12 This lowering of activation energy is encountered often when dealing with radical chain mechanisms.17 The observation in the present experiment leads further to the conclusion that other non-free-radical chain pathways having been eliminated, a radical chain mechanism is operating. The first step is postulated to be C-N bond scissure and the mechanism described by eq 1,2,8,9,10, and 11is therefore taken to be the dominant one for the decomposition of MMA.
References and Notes (1) B. S. Haynes, D. Iverach, and N. Y. Kirov, Symp. (Int.) Combust., [Proc.], 75th (1974). (2) Y. Hidaka, K. Saito, and H. Yamamura, Chem. Lett., 1151 (1973). (3) A. Burcat, G. Skinner, R. Crossley, and K. Scheller, Int. d . Chem. Kinet., 5, 345 (1973). (4) Q.I.Smith and R. F. Sawyer, "The Thermal Pyrobsis of Methylamine", Report No. UCB-ME-76-2, Berkeley, The University of California, April 1976; presented at the Western States SectionlCombustion Institute Paper No. 76-221, Salt Lake City, Utah, 1976. (5) E. Meyer and H. Gg. Wagner, Z . Phys. Chem. (Frankfurt am Maln), 89, 329 (1974). (6) R. E. Weston, Jr., and H. A . Schwarz, "Chemical Kinetics", Prentice-Hall, Englewoad Cliffs, N.J., 1972, p 119. (7) E. A. Dorko, R. W. Crossley, U. Grimm, G. W. Mueller, and K. Scheller, J . Phys. Chem., 77, 143 (1973). (8) D. R. Stull, E. F. Westrum, Jr., and G. C. Swike, "The Chemical Thermodynamics of Organic Compounds", Wiley, 1969, p 461. (9) A. P. Gray and R. C. Lord, J. Chem. Phys., 26, 692 (1957). (10) R. H. Pierson, A. N. Fletcher, and E. St. Clair Gantz, Anal. Chem., 28, 1218 (1956). (1 1) E. A. Dorko, U. Grimm, G. W. Mueller, and K. Scheller, "Shock Tube Isomerization of Cvcloorooane with Infrared Analysis", ARL 7 1-0089. . Wright-Patterson AFB, OH, ARL, May 1971. (12) G. W. Mueller, N. R. Pchelkin, and E. A. Dorko, Program REAaN, submitted to the Quantum Chemistry Program Exchange. Indiana University, Bloomington, IN. (13) P. K. Robinson and K. A. Holbrook, "Unimolecular Reactions", Wiley-Interscience, New York, 1972. (14) J. C. Wert, 111, M.S. Thesis, Air Force Institute of Technology, Wright-Patterson Air Force Base, OH, Dec 1977. (15) P. Gray and J. C. J. Thynne, Trans. Faraday Soc., 59, 2275 (1963). (16) R. C. Weast, Ed., "Handbook of Chemistry and Physics", 55th ed, Chemical Rubber Co. Press, Cleveland, OH, 1974-1975. (17) A. A. Frost and R. G. Pearson, "Kinetics and Mechanism", 2nd ed, Wiley, New York, 1953.
