T H E
J O U R N A L
OF
PHYSICAL CHEMISTRY Registered in U.S. Patent Office 0 Copyright, 1978, by the American Chemical Society
VOLUME 82, NUMBER 15
JULY 27,1978
Kinetics and Mechanism of the Thermal Decomposition of Ethyl Cyanide Keith D. King" and Richard D. Goddard Department of Chemical Engineering, University of Adelaide, Adelaide, South Australia 500 1 (Received March 6, 1978)
The pyrolysis of ethyl cyanide in the presence and absence of radical-chain inhibitors has been investigated in the temperature range 896-1020 K using a stirred-flow reactor at atmospheric pressure with nitrogen as the carrier gas. In the absence of additives the decomposition proceeds by a homogeneous free-radical chain process to yield HCN, C2H4, H2,and C2H3CNas major products. On the assumption of first-order kinetics rate constants for the overall reactions C2H5CN HCN + C2H4 (3) and C2H5CN H2+ C2H3CN(4) are given by the Arrhenius expressions log (k3/s-') = (14.3 f 0.4) - (283.7 f 6.7) kJ mol-'/2,3RT and log (k4/s-l) = (12.1 f 0.3) - (247.3 f 5.9) kJ mol-'/2.3RT. Some aspects of a radical-chain mechanism are discussed. Addition of large amounts of PhNH2 or PhCH, relative to ethyl cyanide completely suppresses the overall reactions 3 and 4 and the mechanism of ethyl cyanide decomposition becomes C2H5CN CH3 + CH2CN (11, CH, + RH CHI + R, CH2CN + RH CH3CN + R where RH is PhNH2or PhCH,. In the presence of excess PhNHz the first-order rate constants for CH&N formation, which are indicative of the unimolecular C-C bond fission reaction 1, are reasonably consistent with the Arrhenius expression log (k1/s-') = 15.25- 338.9 kJ mo1-'/2.3RT. The results are in good agreement with predictions based on the very low-pressure pyrolysis (VLPP) of alkyl cyanides and provide further confirmation that the stabilization energy of the cyanomethyl radical is -21 kJ mol-'.
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Introduction Recently we have shown from very low-pressure pyrolysis (VLPP)experiments that isopropyl cyanide,' n-propyl cyanide,2 isobutyl cyanide,2 and tert-butyl cyanide3 decompose mainly (>go%) via C-C fission adjacent to the CN group. HCN elimination is a minor process. These studies have yielded values for the heats of formation and stabilization energies of a-cyanoalkyl radicals. The results have been used to predict that the Arrhenius expressions for the decomposition of ethyl cyanide (EtCN) via the unimolecular pathways C2H6CN CH3 + CH2CN (1) and CzHSCN C2H4 + HCN (2) are given by2 log (kl/s-') = 15.4 - 338.9 kJ molF1/2.3RT and log (k2/s-l) = 13.6 - 343.0 kJ mol-'/2.3RT. Previous work on the pyrolysis of EtCN has been carried out by Hunt, Kerr, and Trotman-Dickenson (HKT); and Dastoor and Emovon (DE).5 HKT, using the aniline-
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0022-3654/78/2082-1675$01 .OO/O
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carrier technique over the temperature range 959-1038 K, obtained results indicative of C-C fission (log (h/s-l) = 14.1 - 304.2 kJ mol-'/2.3RT), C2H4formation (log (h/s-') = 15.0 - 323.4 kJ rnol-'/2.3RT), and H2 formation (log (h/s-l) = 12.4 - 270.3 kJ m01-~/2,3RT),the latter two being ascribed to molecular processes. In marked contrast, DE, using a conventional plug-flow system over the temperature range 803-943 K, claim to have observed HCN elimination only, with Arrhenius parameters given by log (k/s-l) = 13.11 290.7 kJ mol-'/2,3RT. The parameters for C-C fission and HCN elimination obtained by H K T and DE differ markedly from the estimates based on the VLPP studies of alkyl cyanides. The A factor for HCN elimination obtained by H K T is much too high for a reaction of this type6 and the bond fission activation energy is quite inconsistent with the now well-established value of -21 kJ mol-I for the radical stabilization energy "of the CN g r ~ u p . ' - ~ ~Furthermore, ~-'~ as suggested by a recent study and analysis of H2elimination reactions,ll the unimolecular elimination of H 2 from EtCN a t the rates observed by H K T is highly unlikely. It is surprising that DE did not 0 1978 American
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The Journal of Physical Chemistry, Vol. 82, No. 15, 1978
observe C-C fission unless, of course, HCN and C2H4 were formed by a radical chain process initiated by C-C fission. Chain decomposition of EtCN was suggested by O'Neal and Benson12 when they made an analysis of the data of H K T and we have shown2J3that free radical chain processes make a significant contribution to HCN formation in the pyrolysis of isopropyl cyanide under experimental conditions similar to those of H K T and DE. Chain decomposition of EtCN under their experimental conditions seems very likely. In an attempt to resolve some of the discrepancies, and as part of our continuing studies of the thermal reactions of organic cyanides, we have carried out a reinvestigation of the thermal decomposition of EtCN using a stirred-flow reactor t e ~ h n i q u e .We ~ have already shown2J4that EtCN is thermally stable under VLPP conditions at temperatures up to ca. 1150 K.
Experimental Section The apparatus and experimental procedure have been described in detail p r e v i o ~ s l y .The ~ flow system, which was of conventional design, incorporates the general techniques of Herndon and co-workers15 and the stirred-flow reactor design of Mulcahy and Williams.16 The apparatus was operated as a forced-flow system at around 1-atm pressure using high-purity nitrogen as a carrier gas. The reactant concentration in the gas stream was ca. mol dm-3. The reaction vessel, which was constructed from Vycor, was spherical with a volume of 179.3 cm3 and a surface-to-volume ratio of 0.9 cm-', which could be increased to 2.7 cm-l by the addition of pieces of quartz tubing. Analysis and identification of products were achieved with a Shimadzu GC-4A gas chromatograph fitted with flame ionization and thermal conductivity detectors. Three columns were used: (a) 10% w/w diethylene glycol adiapate on 60-80 mesh Gas-Chrom Z, (b) 10% w/w squalane on 60-80 mesh Gas-Chrom Z, and (c) Porapak Q + R. The first two columns were used for HCN and organic cyanides and the latter was used for H2 and low-molecular weight hydrocarbons. Calibration mixtures of various products with the reactant were prepared. EtCN (K & K) was degassed and vacuum distilled bulb-to-bulb. Its purity was checked by gas chromatography and mass spectrometry. In addition, authentic samples were used for product identification and gas chromatographic calibrations. These were vinyl cyanide (Koch-Light), methyl cyanide (Ajax, Univar), CHI (Matheson, C.P.), C2H4 (Matheson, C.P.) and H2 (J. T. Baker, high purity). HCN was prepared by the addition of aqueous HzS04to aqueous KCN. The gas was collected a t 0 "C, dried, and distilled bulb-to-bulb in vacuo until pure (gas chromatography and mass spectrometry). Aniline and toluene, for use as inhibitors, were high purity analytical grade. Results and Discussion Preliminary experiments showed that in the absence of inhibitors the main products of the pyrolysis of EtCN over the temperature range 896-1020 K were H2, C2H4, HCN, and C2H3CN. A t the highest temperature small amounts of CHI and CH3CN were observed. In the presence of aniline (PhNHJ or toluene (PhCH3) the formation of all main products showed a dramatic decrease but the formation of CH4 and CH3CN increased. These observations indicate that EtCN decomposes predominantly via a radical chain mechanism initiated by C-C fission. In the presence of inhibitor, the radicals CH3 and CHzCN formed in the initial step, are scavenged.
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K. D. King and R. D. Goddard
C2H5CN CH3 + CH,CN (1) CH3 + PhNH2 (PhCH3) CHI + P h N H (PhCH,) CHzCN + PhNH2 (PhCH3) CH3CN + PhNH (PhCH,) PhNH2 and PhCH3, which are used in the aniline- and toluene-carrier techniques, are known effective radical traps provided the temperature is not too high.12 Kinetic measurements were carried out first of all for the uninhibited decomposition. On the basis of the stoichiometry CzH5CN = C2H4 HCN (3) CzHSCN = C2H3CN H2 (4) first-order rate constants may be calculated from the expressions h3 = (U/V)([HCNl /[EtCNl) and h4 = (V/V)([CzH,CNl/[EtCNI)
+
+
where U is the volumetric flow rate a t the reaction temperature, V is the reactor volume, and [HCN], [C2H3CN],and [EtCN] are the steady-state concentrations of product and reactant species. Experimental data were obtained over the temperature range 898-1020 K for residence times of 3.8-7.7 s. The data are not sufficient to firmly establish the reaction order but we have proceeded on the basis of first-order kinetics in order to have a fair comparison with the results of HKT and DE. The results are summarized in Table I. A least-squares treatment of the data yields the Arrhenius expressions log (k3,'s-l) = (14.3 f 0.4) (283.7 f 6.7) kJ mol-'/2.3RT and log ( k 4 / ~ - ' ) = (12.1 f 0.3) (247.3 f 5.9) kJ m01-~/2,3RT where the error limits correspond to one standard deviation. The Arrhenius plots are shown in Figure l. Rate constants measured with the packed reaction vessel showed no significant deviation from the Arrhenius plots; the data points are shown in Figure 1. Some quantitative studies were carried out on the effect of the inhibitors PhCH, and PhNH2. The former was used by DE and the latter by HKT. A series of experiments was carried out a t temperatures of 900, 941, and 990 K, with [PhCH3]o/[EtCN]oranging up to -60; the two lower temperatures are within the range of D E S experiments and the highest temperature falls within the range of HKT's experiments. Gas chromatographic analysis using the Porapak Q + R column showed a rapid depletion of H2 and CzH4 formation with increasing inhibitor concentration. Due to the similar retention times of EtCN and PhCH3 on the diethylene glycol adiapate column, the squalane column was used but this allowed only the formation of HCN to be followed, not CH3CN or C2H3CNdue to their similar retention times on this column. First-order rate constants for HCN formation were determined for different concentrations of inhibitor. The results are summarized in Table 11. It can be seen that the rate constant declined by a measurable factor of -300 over the range of [PhCH,] used. Also shown are the rate constants calculated from the Arrhenius expressions of H K T and DE. At all three temperatures their values are higher than the lowest rate constants determined here, which is a fairly clear indication that radical chain processes were occurring
The Journal of Physical Chemistry, Vol. 82, No. 15, 1978
Thermal Decomposition of Ethyl Cyanide
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TABLE I: Stirred-Flow Reactor Results for the Uninhibited Pyrolysis of Ethyl Cyanide 10*k/s- *
T,K
flow rate, cm3 s-I
[ HCN ] / [ EtCN ]
[ VCN ] / [EtCN]
HCN
VCN
888.6 890.3 896.2 895.5 895.7 914.4 915.6 921.5 932.4 934.0 934.2 938.2 950.2 954.2 954.7 954.8 974.1 974.5 974.6 991.7 994.1 994.1 1011.6 1019.7 1020.2
21.8 30.7 26.5 32.7 23.3 33.9 24.2 22.2 25.5 34.6 42.5 32.4 23.3 44.0 34.9 26.1 45.5 36.0 26.0 26.1 36.8 46.1 46.9 37.8 27.3
0.0353 0.0146 0.0379 0.0318 0.0453 0.0740 0.0689 0.169 0.156 0.130 0.123 0.238 0.431 0.248 0.344 0.426 0.461 0.558 0.738 1.34 1.20 0.938 1.45 2.67 3.61
0.0230 0.0187 0.0237 0.0197 0.0446 0.0597 0.0830 0.0887 0.131 0.108 0.0950 0.0996 0.231 0.159 0.21 0 0.268 0.291 0.359 0.455 0.687 0.605 0.494 0.730 1.16 1.35
0.43 0.25 0.56 0.58 0.59 1.40 0.93 2.10 2.21 2.5 2.9 4.3 5.6 6.1 6.7 6.2 11.7 11.2 10.7 19.5 24.6 24.1 37.9 56.4 55.0
0.28 0.32 0.35 0.36 0.58 1.13 1.12 1.10 1.86 2.09 2.25 1.80 3.0 3.9 4.1 3.9 7.4 7.2 6.6 10.0 12.4 12.7 19.1 24.5 20.6
~~~
IO
1
1
lo-’
lo-’
-
. I
v)
Y
.lo-‘
IO-^
x
IO-^
lo-*
10-4
0.95
1.o
105
lo3 KIT
IO-^ 090095
100
1.05
110
115
120
lo3 KIT
Figure 1. Arrhenius plot for the uninhibited pyrolysis of ethyl cyanide: (0)CzH,CN formation, log (k4/s-’) = 12.1 - 247.3 kJ mor1/2.3RT; (0) HCN formation, log ( k 3 / s - ’ )= 14.3 - 283.7 kJ mol-’/2.3RT; (0 and U) packed reactor. Numerals indicate overlapping points.
under their experimental conditions. Surprisingly, DE found that the rate of formation of HCN was not affected by the presence of PhCH, in approximately equal amounts to EtCN. The rate constants calculated from our predicted Arrhenius expression are considerably below all measurements. Thus while the results of these experiments do not provide confirmation of our predicted Arrhenius parameters they do show that the unimolecular elimination of HCN must be substantially slower than reported by either H K T or DE. A series of experiments was carried out over the temperature range 934-1075 K with PhNHz as inhibitor. At [PhNHz]o/[EtCN]o N 4,the first-order rate constants for HCN formation and C2H3CNformation were a factor of
Figure 2. Arrhenius plot of apparent first-order rate constants for C-C bond fission in ethyl cyanide. [PhNHz]o/[EtCN]o = R ; R = 4 (@); R = 6 (0);R 8 (0);ref 4 results (A). The solid line corresponds to log (k1/s-’)= 15.4 - 338.9 kJ rnor1/2.3RTand the broken line to log (k’1s-I) = 15.25 - 338.9 kJ mo1-’/2.3RT.
-20 less than the values in the absence of inhibitor but were reasonably close to the results of H K T for C2H4 formation and H2 formation, respectively. H K T did not indicate the value of the ratio [PhNH2Io/[EtCNloused by them. First-order rate constants for CH3CN formation were similar to those obtained by HKT for CHI formation. However, on increasing [PhNH2lo/[EtCNIoto 6 the rate constants for CH3CN declined by a factor of ca. 2-3. At [PhNH,],/[EtCN], N 8, the rate constants were approximately the same. A t [PhNH2]o/[EtCN]o> 8 the dominating presence of PhNH2 interfered with gas chromatographic analysis of the reaction products. If all CH2CN and CH3 radicals are scavenged by PhNHz then the first-order rate constants for CH3CN (or CHI) formation are a direct measure of reaction 1. The experimental rate constants are plotted in Figure 2. At the high end of the temperature range the rate constants fall away,
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K. D. King and R. D. Goddard
of Physical Chemistry, Vol. 82, No. 15, 1978
TABLE 11: Effect of PhCH, on the Formation of HCN from Ethyl Cyanide Pyrolysis 900 K [PhCH, I,/ [EtCNla 0 1.4 5.5 20 > 20
HKTa DEb VLPPC a
941 K 103k/s-1
9.7 2.3
0.4 E0.1 < 0.08
0.17 0.174 5 x 10-4
990 K
[PhCH, 1 a I [EtCN 1 0 0 1.0 1.2 2.2 8 19
103 his- 1 33.0 14.0 3.2 0.9 0.5 -0.1
[PhCH, 1 0 I [EtCNI a 0 1.2 3.5 8.5 35 60 -110
290 85 35 11 3.2 1.0 -0.3
HKT DE VLPP
1.11 0.94 4 x 10-3
HKT DE VLPP
8.7 5.9 5x
103 his - 1
Hunt, Kerr, and Trotman-Dickensonp log ( k / s - ' )= 15.0 - 323.4 kJ mol-'/2,3RT. Dastoor and E m ~ v o n ,log ~ (kls-') Predicted from VLPP studies,Zlog ( k l s - ' ) = 13.6 - 343.1 kJ mo1-'/2.3RT.
= 13.11 - 290.7 kJ mol-'/2.3RT.
probably due to decomposition of CH,CN.17 We did not measure quantitatively CH4 formation. Also shown in Figure 2 is the line corresponding to our predicted Arrhenius parameters for c-C fission, ise., log (k1/s-I) = 15.4 - 338.9 kJ m01-~/2,3RT. The present results, while insufficient for a proper least-squares treatment, do support this expression. Indeed, if the data points a t the highest temperatures are discarded for the reason given above, then the results for [PhNHz]o/[EtCN]o 6 and N 8 are consistent with the expression log (k1/s-I) = 15.25 - 338.9 kJ m01-~/2,3RT.Thjs means that AHf0,,&H2CN) = 244.8 kJ mol-l and SEo(CHzCN) N 21 kJ mol-I determined previously2are confirmed. Also the A factor certainly gives us added confidence in the estimated A factors used in our VLPP studies of alkyl ~ y a n i d e s . l - ~ The experimental evidence which we have obtained clearly shows that the pyrolysis of EtCN under conventional conditions proceeds via a radical-chain process initiated by C-C bond fission. It is also clear that this chain process could not have been completely suppressed under the experimental conditions of H K T and DE. As mentioned previously,2 a plausible radical-chain scheme of initiation and propagation is as follows: CzH5CN CH, + CHzCN (1) CH, + C2H6CN CH4 + CH,CHCN (4) CH2CN C2H5CN CH,CN + CH,CHCN (5) CH,CHCN CzH3CN + H (6) H CzH5CN HCN CH3CHz (7) H C2H5CN H z + CH,CHCN (8) CHSCH2 C2H4 + H (9) Reactions of the type CH, CzH5CN CH4 CHzCHzCN (4a) CH2CN CzH5CN CH,CN + CHzCHzCN (5a) have not been included because we expect Ela > E4 and ESa> E5 by -21 k J mol-' in accordance with the effect of CN substitution on metathetical reactions.l0!l8 For approximately equal A factors, this difference amounts to a factor of 12 in rate constant a t 1000 K. Even if there was some contribution from (4a) and (sa) the radical CH2CH2CNshould decompose entirely by the reaction CHzCHzCN C2H3CN + H (64 which yields exactly the same products as reaction 6. The alternative decomposition pathway CHzCH2CN CzH4+ CN (6b) cannot compete if E6b- E 6 a is approximately equal to the difference in endothermicities, which may be calculated
+
+ +
+ +
-----
-
+
+
TABLE 111: Arrhenius Parametersa for Radical-Chain Steps reaction 1 6 7 8 9 a A in s - ' or
log A 15.4 13.8 10.7 10.8 13.5 dm3 mol-'
E
ref
338.9 193.7 17.6 27.6 170.3
2 est, 11 est, 17, 20 est, 17 11
s-l ; E in kJ mol-'.
from thermochemical datalJ to be -90 kJ mol-l. This holds because the reverse, addition reactions have similar activation energies.l9 Assuming steady state and long chains, the relative concentration of chain propagating radicals are given by [CH,CHCN]/[H] = k8[EtCN]/h6 and [CH,CH,]/[H] = h7[EtCN]/h9 The Arrhenius parameters for the relevant radical-chain steps are listed in Table 111. The parameters for reaction 9 are well known.ll The parameters for reaction 6 have been estimated from those for the analogous reaction of the isopropyl radicalll by correcting A for reaction path degeneracy and E for cyano stabilization. Similar principles were applied for the estimation of the parameters for reaction 8, using the parameters for the analogous reaction of H atoms with propanel' as a basis. Several workersz0 have shown that H atoms are capable of pseudo-CN abstraction from both MeCN and EtCN a t a faster rate than H abstraction; the parameters for reaction 7 have been estimated from those for reaction 8 and relative rate data.z0 Thus with [EtCN] = lo-, mol dm-, and T , = 955 K, the calculated relative concentrations of radicals are [CH3CHCN]/[H] N 1200 and [CH3CHz]/[H] N 90, Now, if the initiation step is first order under the experimental conditions (confirmed by an RRK calculation) then termination steps of the pp or ppM type,21i.e. H CH,CHCN products (10) H + CHBCHz products (11) H H + M - Hz + M (12) will give rise to first-order kinetics for the overall decomposition reactions. However, because H atoms are in least concentration by a substantial factor then pp terminations, i.e. 2CH3CHCN products (13) (14) 2CH3CH2 products CH3CHz CH,CHCN products (15) are more likely. These will give rise to one-half-order
+
+
+
--
-- -
Ion-Molecule Reactions in CH3CI- and C,H,CI-NH,
kinetics for first-order initiation. More than likely chain termination, if homogeneous, is some complex mixture of the above possibilities. Clearly, more detailed experimental studies are needed for the elucidation of the complete mechanism of the uninhibited decomposition. In this regard, a conventional static system should prove more fruitful than the present apparatus. However, the main aims of this work. have been achieved. T h a t is, we have shown that the thermal decomposition of EtCN is essentially a radical-chain reaction and the unimolecular elimination of HCN and H2 do not appear to occur a t rates indicated by the results of other workers; we have shown also that the first-order rate constant for reaction 1is very close to predictions based on our previous VLPP studies of alkyl cyanides which, in essence, confirms the value of e 2 1 kJ mol-l for the stabilization energy of the cyanomethyl radical. Achnowledgment. R.D.G. acknowledges the receipt of a Commonwealth Post-Graduate Research Award.
References and Notes (1) K. D. King and R. D. Goddard, J. Am. Chem. Soc., 97, 4504 (1975). (2) K. D. King and R. D. Goddard, Int. J. Chem. Kinet., 7, 837 (1975); E for C-C fission in EtCN is 334.7 kJ mol-' at 300 K and 338.9 kJ mol-' at 1100 K.
The Journal of Physical Chemistry, Vol. 82, No. 15, 1978
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K. D. King and R. D. Goddard, J. Phys. Chem., 80, 546 (1976). M. Hunt, J. A. Kerr, and A. F. Trotman-Dickenson, J . Chem. Soc., 5874 (1965). P. N. Dastoor and E. U. Emovon, Can. J . Chem., 51, 388 (1973). H. E. O'Neal and S. W. Benson, J. fhys. Chem., 71, 2903 (1967). D. A. Luckrafl a d P. J. Robinson, Int. J. Chem. Kinet., 5, 137 (1973). K. D. King and R. D. Goddard, Int. J . Chem. Kinet., 7, 109 (1975). K. D. King and R. D. Goddard, Int. J . Chem. Klnet., in press. Y. Gonen, L. A. Rajbenbach, and A. Horowitz, Int. J. Chem. Kinet., 9, 361 (1977). 2.B. Alfassi, D. M. Golden, and S. W. Benson, Int. J. Chem. Klnet., 5, 991 (1973). S. W. Benson and H. E. O'Neal, Natl. Stand. Ref. Data Ser., Natl. Bur. Stand., No. 21 (1970). K. D. King and R. D. Goddard, unpublished work. K. D. King, unpublished work. W. C. Herndon and L. L. Lowry, J. Am. Chem. Soc., 88, 1922 (1964); W. C. Herndon, M. B. Henley, and J. M. Sullivan, J . fhys. Chem., 67, 2842 (1963). M. F. R. Mulcahy and D. J. Williams, Aust. J. Chem., 14, 534 (1961). T. W. Asmus and T. J. Houser, J . fhys. Chem., 73, 2555 (1969). Y. Gonen, A. Horowitz, and L. A. Rajbenbach, J. Chem. Soc., Farachy Trans. 1 , 72, 901 (1976). J. A. Kerr and M. J. Parsonage, "Evaluated Kinetic Data on Gas Phase Addition Reactions", Butterworths, London, 1972; G. E. Bullock and R. Cooper, Trans. Faraday Soc., 67, 3285 (1971). W. Forst and C. A. Winkler, Can. J . Chem., 33, 1814 (1955); D. E. McElcheran, M. H. J. Wijnen, and E. W. R. Steacie, ibid., 38, 321 (1958); J. W. S. Jamieson, G. R. Brown, and J. S. Tanner, ibid., 48, 3619 (1970). K. J. Laidler, "Reaction Kinetics", Vol. I, Pergamon Press, London, 1963.
Positive Ion-Molecule Reactions in the Methyl Chlorideand Ethyl Chloride-Ammonia Systems Zygmunt Luczynskl' and Jan A. Herman* Centre de Recherches sur les Atomes et les Mol6cules et D6parlement de Chimie, Universit6 Laval, Qusbec, GIK 7 f 4 , Canada (Received February 6, 1978) Publication costs assisted by the Universit6 La Val
The ion-molecule reactions of ions formed by photoionization in the 11.6-11.8-eV region in CH3C1-NH3 and C2H5C1-NH3binary systems were investigated by high-pressure mass spectrometry (0.01-1 Torr). The dialkylchloroniumions formed at higher pressures are very effectively transformed in the presence of small quantities of ammonia into solvated alkylammonium ions [RNH3+(NH3),],where R = CH3 or C2H5 and m I3. In the presence of high concentrations of ammonia, practically the only observed positive ions in both binary systems are solvated ammonium ions, [NH4+(NH3),],where n I 4.
Introduction Ionic processes in gaseous methyl and ethyl chlorides have been intensively investigated in the past decade by various mass spectrometric techniques,2-10in conventional radi~lysis,~ and in vacuum ultraviolet photolysis.ll Results obtained by mass spectrometry allowed kinetic data and a kinetic scheme to be determined rather precisely for ion-molecule reactions in both pure compounds and their mixtures with hydrogen and nitrogen. Most of the positive parent ions formed in irradiated or photoionized CH&l or C2H6C1react with the parent molecules to give CH3C1H+ and C2H5C1H+species, which in turn condense very efficiently with a second parent molecule with HC1 elimination, thus forming dialkylchloronium ions, [ (R),Cl]+. These ions and their solvates, [(R),Cl+(RCl),], are unreactive toward CH3C1and C2H5C1neutrals, and, therefore, they seem to be the end-product positive ions, their yields being almost 100% a t a pressure of 1 Torr and moderate temperatures. It is interesting for the radiation chemistry 0022-365417812082-1679$0 1.OO/O
of gases to investigate the modifications of the described ionic processes in pure chloroalkanes in the presence of scavengers. The interpretation of results obtained by the scavenging-of-positive-ions technique generally used in radiation chemistry may be difficult in view of the unknown processes occurring between ions and the scavenger molecules. However, the direct observation of the ionic process sequence is feasible in a high-pressure mass spectrometer, and some assumptions made in stationary radiolysis experiments can now be confirmed or disproved. The object of the present study is to obtain information on positive ion-molecule reactions in CH&1 and C2HSC1 in the presence of NH3, which is a typical positive ion scavenger.
Experimental Section The measurements were done in the photoionization high-pressure mass spectrometer already described, where some minor modifications were incorporated.12 In order 0 1978 American
Chemical Society
