J. Phys. Chem. 1985, 89, 3109-31 13
3109
Thermal Decomposition of Formaldehyde at High Temperatures KO Saito,* Terumitsu Kakumoto, Yoshihiro Nakanishi, and Akira Imamura Department of Chemistry, Faculty of Science, Hiroshima University, Naka- ku, Hiroshima 730, Japan (Received: February 14, 1984; In Final Form: March 19, 1985) The thermal decomposition of formaldehyde was investigated behind reflected shock waves by monitoring time-dependent C H 2 0 and CO concentrations by IR emission and H-atom concentration by the ARAS method. From the IR emission experiment, it was found that the C H 2 0 decay and the CO formation rates were the same and showed a strong dependence on the reactant mole fraction. An Arrhenius expression for the second-order rate constant of CHzO + Ar CHO + H Ar was obtained as k l = 1015.s0 exp(-75.0 kcal mol-'/RT) cm3 mol-' s-I over the temperature range 2200-2650 K. Rate constants determined from the rate of H-atom production agreed with this expression. The alternative channel C H 2 0 + AI CO + H2 + Ar appeared to have a smaller rate under the experimental conditions studied. Relative rates of the competing channels are discussed on the basis of recent ab initio calculations for their threshold energies.
+
-
-
Introduction Formaldehyde plays an important role as an intermediate in the oxidation of hydrocarbons, where it reacts in a complicated manner.' The mechanism of its formation and decomposition is therefore an important subject for clarification of the combustion mechanisms of hydrocarbons. The initial decomposition step has receive a great deal of attention in theoretical and experimental chemistry. However, there still remains a controversy about the rates of the molecular elimination and radical formation reactions in thermal systems. There are several studies of C H 2 0 thermal decomposition at high temperatures using shock tube techniques.2d Most of them were done by monitoring the C-H stretch IR emission intensity as a function of time.3-5 It was generally found that the decomposition rate depended strongly on reactant concentration, showing that the reaction proceeded by a chain mechanism. From analyses of the apparent decay data considering the following reactions, essentially the same rate constants were obtained for reaction 1 by all authors. CHzO Ar C H O H Ar (1)
+
CH20
+ -
--
H
+ Ar CHO + H
CHO
+ + C H O + H2 CO + H + Ar C O + H2
(2) (3)
(4)
Just and co-woikers investigated the decomposition rates for highly dilute mixtures (down to 2 ppm in Ar) using the ARAS method to follow the H-atom concentration as a function of time.6 They also found similar values of k l . However, contrary to other results, they found a larger rate constant for the competing channel C H 2 0 Ar H2 C O Ar (la)
+
-
+
+
According to their results, k,, is larger than kl by a factor of 9 mol ~ m - ~ . a t 2000 K and a total density of ca. 1 X Using a static system, Batt et al.' studied the decomposition over a temperature range of 750-850 K in mixtures diluted by methane ( [CH20]/ [CHI] = 10-3)at total pressures of 700-800 torr. They examined the participation of channel l a using CDIO (1) For example: Peters, J.; Mahnen, G. 'Proceedings of the 14th Symposium (International) on Combustion"; The Combustion Institute: Pittsburgh, 1973; p 133. Boni, A. A,; Penner, R. C. Combust. Sci. Technol. 1977, 15, 99. (2) Gay, I. E.;Glass, G. P.; Kistiakowsky, G. B.; Niki, H. J. Chem. Phys. 1965, 43, 4017. (3) Schecker, H. G.; Jost, W. Ber. Eunsenges. Phys. Chem. 1969,73,521. (4) Dean, A. M.; Craig, B. L.; Johnson, R. L.; Shultz, M. C.; Wang, E. E. 'Proceedings of the 17th Symposium (International) on Combustion"; The Combustion Institute: Pittsburgh, 1979; p 577. ( 5 ) Dean, A. M.; Johnson, R. L.; Steiner, D. C. Combust. Flame 1980, 37, 41. (6) Just, T. In 'Shock Waves in Chemistry"; Lifshitz, A,, Ed.; Marcel Dekker: New York, 1981; Chapter 6. (7) Batt, L.; Alvaradc-Salinas, G.; Reid, J. A. B.; Robirison, C.; Smith, D. B. 'Proceedings of the 19th Symposium (International) on Combustion"; The Combustion Institute: Pittsburgh, 1982; p 8 1.
0022-3654/85/2089-3109$01.50/0
and found that k l awas negligibly small compared with k , under their experimental conditions. The two-channel dissociation mechanism has been investigated by many photochemical and photophysical researcher^.^-'^ It was found that the molecular channel l a occurred by excitation to a level above the SI band origin (80.6 kcal mol-' above the So). The threshold energy for radical formation (1) was found to be in the range of 85-87 kcal mol-', within experimental error of the best thermodynamic values. It was suggested that there existed a long-lived metastable form of C H 2 0 between the S1 and the products. Theoretical calculations for the dissociation have been done by several investigator^.'^-'^ The energy barrier for molecular elimination was calculated to be 80 f 2 kcal mol-' with large-scale C I calculations using a large basis et.'^-'^ This value is lower than the heat of reaction for channel 1, 88.7 kcal mol-'. Forst20 calculated the pressure dependence of the rate constants for the two channels considering the tunneling. According to his results, kla becomes important at lower pressures as the result of the tunneling which makes only channel l a faster. From the previous results referred above, it is likely that the rate constant for radical formation ( k , ) has been established at these high temperatures of 150&3000 K within an error of a factor 2. In contrast to this, however, there still remain arguments against accepting the results which demonstrate that the molecular elimination, channel l a , exceeds the radical formation. Clearly, in many studies exact determinations of the initiation rate constants are limited by complicated radical reactions even in systems highly dilute mixtures. This seems to have led to ambiguous conclusions for the pending question about the competing channels. In the present study, therefore, we made an effort especially to evaluate the difference of the rate constants between the two initiation channels in the €herma1 system diluted by Ar. Experimental Section To evaluate the initiation reaction rates, two different observing techniques were employed to follow the time-dependent concentrations of the reactant and products in the shock-tube equipment. (8) Houston, P. L.; Moore, C. B. J . Chem. Phys. 1976, 65, 757. (9) Clark, J. H.; Moore, C. B.; Nogar, N. S . J . Chem. Phys. 1978, 68, 1264. (10) Moortagat, G. K.; Warneck, P. J . Chem. Phys. 1979, 70, 3639. (11) Walsh, R.; Benson, S. W. J . Am. Chem. SOC.1966, 88, 4580. (12) Tand, C. K. Y.; Fairchild, P. W.; Lee, E. K. C. J . Phys. Chem. 1979, 83, 569. (13) Horowitz, A,; Calvert, J. G. Int. J . Chem. Kinet. 1978, 10, 8 0 5 . (14) Ho, P.; Bamford, D. J.; Buss, R. J.; Lee, Y.T.; Moore, C. B. J . Chem. Phys. 1982, 76, 3630. (15) Jaffe, R. L.; Morokuma, K. J. Chem. Phys. 1976, 64, 4481. (16) Goddard, J. D.; Yamaguchi, Y.; Schaefer, H. F. J. Chem. Phys. 1981, 75, 3459. (17) Adams, G. F.; Bent, G. C.; Bartiett, R. J.; Purvis, G. D. J . Chem. Phys. 1981, 75, 834. (18) Frisch, J. J.; Krishnan, R.; Pople, J. A. J. Phys. Chem. 1981,85, 1467. (19) Dupuis, M.; Lester, W. A.; Lengsfield, B. H.; Liu, B. J . Chem. Phys. 1983, 79, 6167. (20) Forst, W. J . Phys. Chem. 1983, 87, 4489, 5234.
0 1985 American Chemical Society
3110
The Journal of Physical Chemistry, Vol. 89, No. 14, 1985
In experiment I IR radiation from shock-heated C H 2 0 and C O product were observed simultaneously. In experiment I1 the time-resolved H-atom concentration was measured by atomic resonance absorption spectroscopy (ARAS) at 121.6 nm. These two experiments were performed in a shock tube made of stainless steel, which was used in previous studies.,' The test section was 4.52-m long and 9.4 cm in inner diameter. It could be evacuated, after baking, to less than 7 X low7torr by means of a 6-in. diffusion pump. In usual experiments it was evacuated to about 1 X torr before each run. After a run, the gas in the tube was roughly pumped out and then dry N, was injected to atmospheric pressure before changing the diaphgram. With this procedure adsorption of water vapor from room air onto the tube wall was suppressed satisfactorily. The driver sections, 1.8-m long and 8 S cm in inner diameter, was evacuated to about 1 X lo-, torr by means of a rotary pump. Formaldehyde was prepared by heating commercial paraformaldehyde in an oil bath (ca. 120 "C) and purified by passing through a dry ice-acetone trap, and then the monomer was condensed at 77 K. After trap to trap distillations, it was stored in a glass flask as a mixture of about 1% in Ar, which was used for preparation of highly dilute mixtures as required in the experiments. Vacuum-UV absorptions at room temperature was used to determine the CHzO concentration in the highly dilute mixtures. The absorption coefficient (base e) was determined to be 6.02 X lo6 cm2 mol-' at 121.6 f 2 nm. This value was used to check the prepared mixtures. On the basis of the observation that the CH,O concentration was constant to within a few percent over a few days, polymerization was not important for the storage of these mixtures in glass flasks. Also, the effect of adsorption on the shock tube wall was ascertained not to be important by the same method. Experiment I. In this experiment, C H 2 0thermal decomposition was monitored by observing thermal IR emission intensity from C H 2 0and CO simultaneously. IR emission from the shock-heated gas was taken out through a pair of CaF, windows mounted on the tube walls 2-cm upstream from the end plate. The radiation from one of the windows was focused on an InSb detector at 77 K after passing through a slit (1-mm wide) and a band-pass filter at 3.37 f 0.13 pm, the CH stretch of CH20. The same optical arrangement was used at the other side with a filter at 4.63 f 0.05 pm, the CO fundamental band. The detector output signals were fed into waveform digitizers for subsequent data analysis. The response time of the optical-electrical system was about 10 ps, small enough for evaluation of the decomposition rate. Experiment ZZ. A time-resolved H-atom absorption study was performed by adding a vacuum-UV absorption spectrometer to the observation station of the shock tube used in experiment I. The windows in this case were 1-mm MgF,. A vacuum-UV monochromator (Tokyokagaku-MCV20) was used to select 121.6 f 2 nm. The light was detected with a solar blind photomultiplier (HTV-R1459). A lyman CY light source was made by a flow of H e containing about 1% H2 excited by 2450-MHz microwave discharge. Because a simple correlation between measured absorption and H-atom concentration does not exist, calibration measurements using mixtures of H2/Ar and H 2 / N 2 0 / A r were performed over the temperature range from 1500 to 2500 K.6 Especially many data were obtained around 2000 K to construct a full calibration curve over a wide concentration range. A slight, but not negligible, temperature dependence was found for the calibration curves. Results A . ZR Emission. Experiments were performed in the tem-
perature range from 1600 to 2650 K, total concentration range (4-10) X lo4 mol ~ m - and ~ , reactant concentration range 50-104 ppm in Ar. Figure 1 shows typical emission traces for C H and CO observed simultaneously. The C H trace decayed exponentially after an initial rapid rise, immediately following the reflected shock
Saito et al.
-
ISRS
50p s
Im
J time
R.S.
-
Figure 1. Typical emission profiles at 3.4 (upper trace) and 4.6 pm (lower trace). IS and RS denote incident shock front and reflected shock front arrivals, respectively. Conditions: T = 1725 K, total density = 0.98 X mol 0.29% CHIO in Ar.
t
'
, - I
I
100
0 t
I
200
m e I ,us
Figure 2. Logarithmic plot of the emission intensities shown in Figure 1. A straight line corresponds to the steepest exponential decay.
front (RS), and reached almost zero, indicating that no significant concentrations of stable C H compounds are produced at these temperatures. An apparent second-order rate constant k, was determined by the equation k, = (-d In Z/dt)/[M], where Z is the intensity and [MI the total concentration. The C O emission signal increased with time to a steady level I, (Figure 1). This profile was fitted to Z = Z,(1 - e-kf),where k, is given by k , = k/[M]. It was found that there was no systematic difference between k,(CH) and k,(CO). Figure 2 shows a logarithmic plot of the time-dependent intensity for the C H emission and I , - Z for the C O emission, where the intensities were scaled such that the peak intensity of Z(CH) was equal to Z,(CO). The two signals overlap except initially. Due to the time constant (- 10 ps) the analysis of high-temperature data was limited especially for C H profiles. It was found from Figure 2 that there was an induction period before exponential decay occurs, indicating a chain mechanism for this reaction as reported by Schecker and Jost3 and Dean et al.4 At high temperatures, however, this induction period decreased, and eventually it became comparable to or less than the time constant of the detectors. The effects of vibrational relaxation of C O upon the rate constant obtained from CO emission have been considered. In previous work21*22 we ascertained that the carbon monoxide produced from the decomposition of glyoxal and formic acid was vibrationally relaxed; therefore, the rates of C O production determined from IR emission were not affected by the vibrational relaxation of CO in the two systems. _____
(21) Saito, K.; Kakumoto, 1182.
T.; Murakami, I. J . Phys. Chem. 1984, 88,
(22) Saito, K.; Kakumoto, Chem. Phys. 1984, 80, 4989.
T.; Kuroda, H.; Torri, S.; Imamura,
A. J
Thermal Decomposition of Formaldehyde
The Journal of Physical Chemistry, Vol. 89, No. 14, 1985 31 11
TIK
2500
Zoo0
1700
I
I
I
I
'0
/
No.718
100
200
300
I
LW
500
t imelps
Figure 5. Concentration profile of H atom for the experimental run shown in Figure 4. The straight line corresponds to kH = 6.7 X lo6 cm3
mol-l
s-l.
T / K
8 I
I
I
4.0
5.0
6.0
2000 l
I
e
lo4 T - ~I K - ~ Figure 3. Plot of the apparent second-order rate constant k, for various mixtures. k, (dotted line) was evaluated from the data for 50, 100, and loo0 ppm mixtures as explained in the text. The line k l , (J) corresponds to the molecular elimination channel given by Just.6
1
I
No.718
50
I
I
5.5
60
I
104T-'/ K-' RS
time
Figure 4. Typical H-atom absorption profile at 121.6 nm. Conditions: T = 1815 K, total density = 1.01 X lo-' mol ~ m - 2UO ~ , ppm CH20 in
Ar. The same situation is also assumed to prevail in the present system. In fact, a t relatively low temperatures and for higher reactant concentrations the decay rate agreed with the Co production rate as shown in Figure 2. At higher temperatures we could not confirm the coincidence of the rate constants from reactant and CO signals, because the CH emission was controlled by the time constant of the detector system. However, from the temperature dependence of k, the CO production rate appeared to be independent of the vibrational relaxation of CO. The vibrational relaxation time had a much lower temperature dependence compared with the activation energies of kaeZ3 As seen in Figure 3, an Arrhenius plot of k, vs. T I for various diluted mixtures, the apparent rate constant was strongly dependent on the mole fraction of reactant, in agreement with previous result^.^,^ The temperature dependence of k, increased with increasing dilution. At temperatures higher than 2500 K the data for 50 and 100 ppm mixtures agreed very closely, suggesting that the decomposition becomes simple such that the decay rate was governed only by a few initial reactions. B. ARAS. Experiments were performed in the temperature range between 1600 and 2010 K, total density range of (9.5-10.6) X IO" mol ~ m - and ~ , reactant concentration range 50-210 ppm in Ar. Figure 4 shows a typical absorption trace a t 121.6 nm. (23) Borrell, P. In 'Transfer and Storage of Energy by Molecules, 2"; Burnett, G . M., North, A. M., Eds.; Wiley: New York, 1969; p 180.
Figure 6. Arrhenius plot of k , obtained from ARAS method for various mixtures. The solid line is the least-squares expression of the data points. k , ( J ) corresponds to Just expression.6
The absorption begins abruptly at the reflected shock front (RS) and approaches a constant level. From the calibration curves, H-atom concentrations were determined as a function of time from each absorption trace. Figure 5 shows the concentration profile for run No. 7 18 given in Figure 4. The second-order rate constant for the H-atom production, kH,was determined from the early Since it can slope by the equation kH = d[H]/dt/[CH,O],[Ar]. be assumed that for the highly dilute mixtures used here reactions 1 and 3 are dominant at the early stage of the reaction, d[H]/dt = 2kl [CH20][Ar]. Thus, we have directly determined k, values from the observed kHas kl = kH/2. Figure 6 shows an Arrhenius plot of k, ( = k ~ / 2 )for various mixtures. Although data points scatter somewhat widely, an Arrhenius expression was obtained by least-squares fit with an error of about a factor of 2 as
k 1 -- 1015.53 exp(-75 kcal mol-'/RT) cm3 mol-'
s-I
Just's expression for k l , shown as a dashed line in Figure 6, confirms that the two results obtained with the ARAS method are in excellent agreement under similar experimental conditions. Discussion A. Rate Constant for Channel I. It has been found that the thermal decomposition of C H 2 0is dominated by secondary radical reactions.2-5 In the present study this was also demonstrated by the drastic change in the decay rate with reactant mole fraction, especially at lower temperatures (Figure 3). Of many likely bimolecular reactions, only four, (1)-(4), were selected as an
3112 The Journal of Physical Chemistry, Vol. 89, No. 14, 1985
Saito et al.
\
- 1 U I
100
300
200
time
500 us
400
time
Figure 8. Comparison of calculated C O profiles (smoothcurves) with experiments. (A) 0.29% CHzO in Ar, T = 1995 K,[MI = 1.05 X mol/cm3. (B) 0.29% C H 2 0 in Ar, T = 1725 K, [MI = 0.98 X mol/cm3. TABLE 1: Elementary Reactions and Their Rate Constants
1% ( A ) /
I
0
40
I
80
1
I
I
120
160
200
(1) (2) (3) (4) us
reaction C H 2 0 Ar C H O H Ar CH2O H H2 C H O C H O + Ar C O H + Ar CHO H 4 CO H2
+ + +
-. + +
+
+
+
+
cm3 mol-l s-I 15.50 13.96 14.20 14.30
E,/ kcal mol-’ 75.0 7.9 14.7 0
ref this work this worka Dean et aL4 Niki et aLb
Evaluated by BEBO method. * Niki, H.; Daby, E. E.; Weinstock, B. ”Proceedings of the 12th Symposium (International) on Combustion”; The Combustion Institute: Pittsburgh, 1969; p 277.
1.0
-
0.5
-
B
0
I
5 d
8
-
L
0
I
0
I 100
I
I
I
I
200
300
400
500
us
time
Figure 7. Comparison of calculated C O profiles (smooth curves) with experiments. (A) 50 ppm CH,O in Ar. Upper trace: T = 2640 K, [MI = 1.10 X mol/cm3; lower trace: T = 2230 K, [MI = 1.13 X mol/cm3. (B) 100 ppm CHIO in Ar. Upper trace: T = 2570 K, [MI = 1.12 X mol/cm3; lower trace: T = 2325 K, [MI = 1.08 X mol/cm3.
important mechanism at the early stage of the reaction in highly dilute mixtures. Computer simulations were performed with this mechanism t o reproduce the observed C O profiles. Then values of k l were evaluated so as to fit the calculated profile to the observed one under the same conditions. Since different values have been reported in the literature for kZ,kj, and k4.several sets of rate constants were used in the calculations. The calculated curves with different sets did not give significantly different results. Figures 7 and 8 show examples of the comparison between ob-
served and calculated CO profiles under several conditions. For the highly dilute mixtures, 50 and 100 ppm (Figure 7), the calculated curves are in good agreement with the experimental curves when the set of rate constants in Table I was used. In additional calculations the competing initial channel l a was added to the mechanism to examine the effect of this reaction, applying the same value for kl, as k , . As expected the results always showed larger CO production rates than without channel la: For some experimental runs of 1000 ppm mixtures the curve fitting was possible using lager k,, and smaller kl values, but this set did not tit the data of the other mixtures. From this fact it was ascertained that the rate constant k , , did not contribute significantly to the overall reaction if k, took the value generally reported by the previous workers. In this respect the results reported by Just6 differ from those of the other investigators. In Figure 3 the line of k,,(J) of Just et al. is shown for comparison with the overall rate constant k , obtained in this work. It is seen that the data for the overall rate cqnstant are much lower than k,,(J) for mixtures containing less than 1000 ppm CH20. Since in a mixture at infinite dilution secondary reactions would be absent, the overall rate constant can be expressed by kl + k,,, and the value of k,, should not exceed the oyerall rate constant. From the curve fitting without channel l a , k , values were determined for each experimental run with mixtures of 50-1000 ppm. These gave an Arrhenius expression k l = IOl5
exp(-75.0 kcal moI-’/RT) cm3 mol-’ s-]
in the temperature range of 2200-2650 K. This expression is in good agreement with that from the direct measurement by ARAS. This agreement supports the finding that k , , des not exceed kl under the present conditions.
J . Phys. Chem. 1985,89, 3 113-3 117 TIK 3oM)
I'
I -
I
2500 I
I
L
1700
2000 I
I
I
6 IO'T"I
I
I
I
5
1500
7
K-'
Figure 9. Comparison of the reported rate constants for channel 1. IR and ARAS are the present results obtained by the two methods; S, Schecker and Jost;' D1 and D2, Dean et D3, Dean et J, Just
k,(G), apparent rate constant by Gay et k,(S), apparent rate et constant by Schecker and J o ~ t and ; ~ k,,(J), channel la by Just et aL6 For high concentration mixtures, however, the calculated CO curves did not agree with the observed data as shown in Figure 8 for the case of 0.29% mixture. The reason for this large discrepancy is obviously due to assuming the simple mechanism for these mixtures. Therefore, it appears to be necessary to include further secondary reactions in the mechanism to explain all the experimental data. However, if the discussion is limited to highly dilute mixtures, the simple mechanism seems to be enough and the Arrhenius expression for kl evaluated from the 4-reaction mechanism appears to be reasonable. This is in essential agreement with the results of the previous worker^.^" That is, the Arrhenius parameters reported are all in the ranges log A = 15.85 f 0.35 and E, = 78 f 3 kcal/mol. Figure 9 summarizes the published data and the present results for kl in Ar above 1500
K. B. Rate Constant for Channel l a . The problem arising from the competing reactions has received a great deal of theoretical and experimental attention as stated above. In the unimolecular
3113
-
decomposition of formaldehyde the following competing channels CHO H have to be ~ o n s i d e r e d :channel ~~ 1, CHzO Ar + Ar with AHr,o"/kcal mol-' = 88.7, and channel l a , C H 2 0 Ar C O Hz Ar with AHr,o/kcal mol-' = -0.4. Channel 1 is a two-center reaction having a loose transition state (TS). Thus the threshold energy Eo for this channel is thought to be close to the heat of reaction if the centrifugal barrier is neglected. In fact, from photochemical studies9J'J3 the energy barrier for this channel was deduced to be in the range 85-88 kcal mol-'. On the other hand, channel l a is a four-center reaction having a tight TS and is thermochemically neutral. In this type of reaction there is a large energy barrier between reactant and products. According to photochemical s t ~ d i e sthis , ~ channel was found to occur by excitation to the SI band origin (80.6 kcal mol-' above the So). Recent ab initio calculations for the energy barrier of this channel have given 88.9,1688.4,'' 83.3,1879.6,18 and 80.919kcal mol-', depending on the basis set and on the CI method used. (All values are corrected for the zero-point energy). It is clear that the barrier becomes lower when a larger basis set and a larger-scale CI method are employed in the calculations. Consequently, the calculated barrier height seems to lie in the range 80 f 2 kcal mol-'. At present, a thermal rate constant for channel l a is given only in the report of Just6 shown in Figure 9 as k,,(J), where kl,(J) - 1016.'5exp(-71.5 kcal mol-'/RT) cm3 mol-' s-I. This is larger than kl by about 1 order of magnitude at 2000 K. It has been suggested that reaction l a might be in the falloff region at the mol ~ m - however, ~; except for the above total density of 1 X report, there are no experimental data which suggest that k , , is larger than k l . Like the present study, Dean et aL4q5found that k l , was not important under the conditions considered. Results of Batt et ale7at lower temperatures from 750 to 850 K again showed no important role of channel la. According to Forst's calculations20 of falloff curves by considering the tunneling, the activation energy of k,, decreased with decreasing total density, while that of kl did not change. This suggests a substantial lowering of k l relative to k,, at very low total densities. Because all the experimental studies so far were limited to a narrow density region, it is necessary to investigate the reaction kinetics over an expanded density range.
-
+
+
+
+
+
Acknowledgment. We thank Mr. Y. Ueda for experimental assistance, Professor I. Murakami for helpful discussions, and Professor D. Munch for careful reading of the manuscript. Registry No. CH20, 50-00-0. (24) JANAF Thermochemical Tables, Supplement J . Phys. Chem. ReJ
Data 1974, 3, 3 11.
Radical Mechanlsm for the Laser-Induced Explosion of Methyl Isocyanide M. J. Shultz,* Robert E. Tricca, S. L. Berets, Christopher Kostas, and Loretta M. Yam Chemistry Department, Tufts University, Medford, Massachusetts 02155 (Received: September 13, 1984)
The laser-induced isomerization of methyl isocyanide has previously been shown to exhibit a marked pressure dependence; large-scale isomerization occurs when threshold conditions have been exceeded. Recent investigations of this large-scale isomerization have shown that the threshold is not due to a thermal mechanism as previously believed. In this work, the characteristics of the threshold are examined and it is shown that the threshold is a result of a radical channel.
I. Introduction In recent years, there have been numerous investigations of laser-induced chemical reactions.'S2 Although it is well-known ( 1 ) Schulz, P. A.; Sudbo,A. S.; Krajnovich, D. J.; Kwok, H. S.; Shen, Y. R.;Lee, Y.T.Annu. Reu. Phys. Chem. 1979, 30, 379.
0022-3654/85/2089-3113$01.50/0
that laser-induced reactions can occur under completely collisionless condition^,^-^ there have been a number of studies doc(2) Ronn. A. Sci. A m . 1979. 114. (3j Coggiola, M. J.; Schulz, P. A.; Lee, Y. T.; Shen, Y . R. Phys. Reu. Left. 1977, 38, 17.
0 1985 American Chemical Society
