J. Phys. Chem. 1984, 88, 2361-2364
2361
Kinetics of the Gas-Phase Reactions of NO3 Radicals with a Series of Alkanes at 296 f 2 K Roger Atkinson,* Christopher N. Plum, William P. L. Carter, Arthur M. Winer, and James N. Pitts, Jr. Statewide Air Pollution Research Center, University of California. Riverside, California 92521 (Received: July 11, 1983)
Relative rate constants for the gas-phase reactions of the NO3 radical, an important reactive species observed in nighttime atmospheres, with a series of alkanes have been determined at 296 f 2 K. Using a rate constant for the reaction of NO3 cm3 molecule-’ s-’, we obtained the following rate constants (in units of radicals with ethene of (6.1 f 2.6) X cm3 m~lecule-~ d):n-butane, 2.0 f 1.0; n-pentane, 2.4 f 1.2; n-hexane, 3.2 f 1.5; n-heptane, 4.1 f 1.8; n-octane, 5.5 f 2.5; n-nonane, 1.2 f 3.3; isobutane, 2.9 f 1.4; 2,3-dimethylbutane, 12.1 f 5.4; and cyclohexane, 4.0 f 1.9. These data, which are the first to be reported for the reaction of NO3 radicals with alkanes, indicate that these reactions are a minor loss process for alkanes in the troposphere, relative to daytime hydroxyl radical attack, but that they may play a role in the formation of nitric acid, a key constituent of acid deposition, during nighttime hours in polluted atmospheres.
Introduction
The nitrate (NO,) radical has long been recognized as an intermediate species in laboratory and environmental chamber NO,-O,-air ~ysternsl-~ and has recently been identified and measured by long-path spectroscopic techniques in ambient nighttime atmosphere^.^-'^ While a number of laboratory studies have been carried out to investigate the and products16 of the reactions of NO3radicals with organics, to date these studies have involved only the simple alkene^,"*'^*'^ aldehyde^,^','^ aromatic hydrocarbon^,^^-'^ methoxy- and hydroxy-substituted aromatics,I3 and dimethyl ~ u l f i d e . ’ ~Clearly, kinetic and product data are required for a wide variety of classes and structures of organics in order to more fully investigate the reactivity of the NO3 radical with organics and to assess the importance of these reactions in the ambient atmosphere as possible sinks for NO, and organics as well as for the formation of nitric acid, a key component of acid deposition. In this study, as part of a comprehensive investigation of gas-phase NO3 radical reactions, we have determined rate constants for the reaction of NO3 radicals with a series of alkanes at room temperature. (1) Johnston, H. S. J . Am. Chem. SOC.1951, 73, 4542. (2) Niki, H.; Daby, E. E.; Weinstock, B. A h . Chem. Ser. 1972, No 113, 16. (3) Demerjian, K. L.; Kerr, J. A.; Calvert, J. G. Adv. Enuiron. Sci. Technol. 1974, 4, 1. (4) Graham, R. A.; Johnston, H. S. J . Phys. Chem. 1978, 82, 254. (5) Platt, U.; Perner, D.; Winer, A. M.; Harris, G. W.; Pitts, J. N., Jr. Geophys. Res. Lett. 1980, 7, 89. (6) Noxon, J. F.; Norton, R. B.; Marovich, E. Geophys. Res. Lett. 1980, 7 , 125. (7) Platt, U.; Perner, D.; Schroder, J.; Kessler, C.; Toennissen, A. J . Geophys. Res. 1981, 86, 11965. (8) Platt, U.; Perner, D.; Kessler, C. “Composition of the Nonurban Troposphere”, 2nd Symposium, Williamsburg, VA, May 25-28, 1982. (9) Platt, U. F.; Winer, A. M.; Biermann, H. W.; Atkinson, R.; Pitts, J. N., Jr. Environ. Sei. Technol., in press. (10) Winer, A. M.; Pitts, J. N., Jr.; Platt, U. F.; Biermann, H. W. “Direct Measurements of Nitrate Radical Concentrations at Desert Sites in California: Implications for Nitric Acid Formation”, Division of Environmental Chemistry, 185th National Meeting of the American Chemical Society, Seattle, WA, March 1983; American Chemical Society: Washington, DC, pp 76-79. (11) Morris, E. D., Jr.; Niki, H. J . Phys. Chem. 1974, 78, 1337. (12) Japar, S. M.; Niki, H. J . Phys. Chem. 1975, 79, 1629. (13) Carter, W. P. L.; Winer, A. M.; Pitts, J. N., Jr. Emiron. Sci. Technol. 1981, 15, 829. (14) Atkinson, R.; Plum, C. N.; Carter, W. P. L.; Winer, A. M.; Pitts, J. N., Jr. J. Phys. Chem. 1984, 88, 1210. (15) Atkinson. R.; Pitts, J. N., Jr.; Aschmann, S. M. J . Phys. Chem. 1984, 88, 1584. (16) Bandow, H.; O h d a , M.; Akimoto, H. J. Phys. Chem. 1980,84,3604.
0022-3654/84/2088-2361$01.50/0
Experimental Section Two experimental techniques, which have been described in detail previo~sly,’~ were used for the determination of NO, radical reaction rate constants. The technique used for the majority of experiments was a relative rate m e t h ~ d ’and ~ ! ~was ~ based upon monitoring the relative decay rates of a series of organics, including one organic whose NO, radical reaction rate constant was known, in the presence of NO3 radicals generated from the thermal decomposition of N205:17
-
N205
M
NO,
NO2 + NO,
+ NO3
M
N2O5
Since the organics studied here react negligibly with N,05 and NO,, then under the experimental conditions employed the sole chemical loss process for these organics was due to reaction with NO, radicals: NO3 + organic
-
-
products
NO3 + reference organic
(3)
products
(4)
Additionally, small amounts of dilution occurred as a result of the incremental additions of N2O5 to the reactant mixture. In these experiments, the dilution factor at time t , Dt, was measured either by the pressure change within the chamber using an MKS Baratron capacitance manometer or, for the all-Teflon chamber, by the volume change in the chamber. During these experiments, was typically 0.004 (Le., -0.4%) per N,05 the dilution factor, D,, addition. With this correction for dilution, then14
\
’
[
5 k4
1 1
[reference o r g a n i ~ ] , ~ In
[reference organic],
- D,
(1)
where [organic],, and [reference organic],, are the concentrations of the organic and reference organic, respectively, at time to, [organic], and [reference organic] are the corresponding concentrations at time t , and k3 and k4 are the rate constants for reactions 3 and 4, respectively. Hence, plots of [In ([organic],,/ [organic],) - DJ against [In ([reference organic],/ [reference organic],) - D,]should yield straight lines of slope k3/k4 and zero
,
(17) Malko, M. W.; Troe, J. Int. J . Chem. Kine?. 1982, 14, 399
0 1984 American Chemical Society
2362 The Journal of Physical Chemistry, Vol. 88, No. 11, 1984
Atkinson et al.
intercept. With this technique the initial concentrations of the ~ , two organic reactants were -(2-3) X lOI3 molecules ~ m - and or three incremental amounts of N2O5 ( ~ ( 5 - 1 2 )X lOI3 molecules cm-, of N2O5 per addition) were added to the chamber throughout the experiments. In addition to these relative rate measurements, experiments were carried out to determine the rate constants k3 for n-heptane and 2,3-dimethylbutane from the increased rate of decay of N205 in N2O5 in Nz05-NOz-alkane-air mixtures. This technique has been described in detail previously,I4 and this reference should be consulted for details. With this technique, and providing that secondary reactions are negligible (see later), the NzO5 decay rate in the presence of an alkane, d In [N,O,]/dt, is given byI4
2,3- D I MET H Y LBUTANE
+
where K is the equilibrium constant for the NO2 NO, a N 2 0 5 reactions and k5 is the rate constant for the background decay of N,05 in the chamber18 (reaction 5). In these experiments the NzO5
-
loss of N2O5
(5)
initial concentrations of the alkanes were -(5-24) X 1014molecules ~ m - with ~ , initial NO, and N205 concentrations of -(5-10) X I O l 3 and -(2-5) X lOI3 molecules ~ m - ‘respectively. ~, The NO2 concentrations, and hence the [alkane]/ [NO,] ratios, were observed to remain constant (to better than &lo%) throughout the duration of these experiments. Experiments were carried out in two environmental chambers. The majority, including all those involving in situ spectroscopic measurements, were conducted in the SAPRC 5800-L thermostated, Teflon-coated evacuable chamber.19 This chamber was fitted with an in situ multiple-reflection optical system interfaced to a Fourier transform infrared (FT-IR) absorption spectrometer. A 62.9-m path length was used in these experiments, resulting in detection limits of 7 X 10” molecules cm-3 for NzO5 and H N 0 3 and 1.2 X lo1, molecules cm-3 for NO, (i.e., -30 and -50 ppb, respectively). Water vapor concentrations were determined from wet bulb-dry bulb measurements of the purified matrix air19,20 used to fill the chamber and were 5 4 X 10l6molecules cm-3 (5% relative humidity at 298 K). All rate constant determinations in this chamber were carried out at 298 f 1 K and -740-torr total pressure of air. Experiments using the relative rate technique to determine rate constants for ethene, n-heptane, and 2,3-dimethylbutane were carried out in a -6400-L all-Teflon chamber at 294 K and -735-torr total pressure of air, as described previ0us1y.l~ For the experiments in which NzO5 decay rates were used to determine individual NO3 radical reaction rate constants, N2O5 and NOz were quantitatively monitored by FT-IR spectroscopy at a spectral resolution of 1 cm-I. In all experiments, the alkanes were analyzed by gas chromatography with flame ionization detection (GC-FID), using a 20 ft X 0.125 in. stainless-steel (SS) column packed with 5% DC703/C20M on 100/120 mesh AW, DMCS Chromosorb G, operated at 333 K. Ethene was analyzed by GC-FID using a 5 ft X 0.125 in. SS column packed with Porapak N (80/100 mesh), operated at 333 K. N 2 0 5was prepared by the method of Schott and Davidson,zl in which NOz was added to a stream of ozonized oxygen at atmospheric pressure and the Nz05collected at 196 K and purified by vacuum distillation. The N,05, was prepared in this manner, always contained 2 1 5% NO, after injection into the chamber. For the experiments in which N2O5 decay rates were used to determine individual NO, radical reaction rate constants, known pressures of the reactants, as monitored by an MKS Baratron (18) Tuazon,E. C.; Atkinson, R.; Plum, C. N.; Winer, A. M.; Pitts, J. N., Jr. Geophys. Res. Lett. 1983, 10, 953. (19) Winer, A. M.; Graham, R. A.; Doyle, G. J.; Bekowies, P. J.; McAfee, J. M.; Pitts, J. N., Jr. Adu. Enuiron. Sci. Technol. 1980, 10, 461. (20) Doyle, G. J.; Bekowies, P. J.; Winer, A. M.; Pitts, J. N., Jr. Enuiron. Sei. Technol. 1977, 1 1 , 45. (21) Schott, G.; Davidson, N. J . Am. Chem. SOC.1958, 80, 1841.
0.02
0
0.06
0.04
0.00
In( [n-HEPTANE]to/[n-HEPTANE]t)-Dt
Figure 1. Plots of eq I for ethene and 2,3-dimethylbutane,with n-heptane as the reference organic. 0.06r
0
6 ’
,
I
I
I
I
0 0.04 0.08 0.12 0.16 In (C2,3-DIMETHYLBUTANElt~/[2,3-DIMETHYLBUTANEl+)-Dt
Figure 2. Plots of eq I for cyclohexane and isobutane, with 2,3-di-
methylbutane as the reference alkane. capacitance manometer, were added to the chamber from calibrated Pyrex bulbs. For the experiments utilizing the relative rate technique, the reactants were introduced into the chamber from all-glass, gas-tight syringes or from Pyrex bulbs by a stream of ultrahigh purity N2. Gas chromatographic analyses of the alkanes and ethene used in this study showed in almost all cases no observable (50.01%) impurities, although n-butane was observed to have -0.007% of trans-2-butene and 0.006% of cis-2-butene impurities. While these alkene impurities have no effect on the relative rate technique, their presence may affect the results of the NZO5 decay rate technique (see Discussion section).
Results In order to avoid interferences in the gas chromatographic analyses employed for the relative rate constant technique, the following sets of organics were studied separately: ethene and n-heptane; the n-alkanes n-butane through n-octane; 2,3-dimethylbutane, isobutane, and cyclohexane; and ethene, n-heptane, and 2,3-dimethylbutane. Representative plots eq I are shown in Figures 1 and 2, and the rate constant ratios kJk4 obtained by least-squares analyses of these data are given in Table I. The error limits given in Table I reflect the two least-squares standard
The Journal of Physical Chemistry, Vol. 88, No. 11, 1984 2363
Gas-Phase Reactions of NO, Radicals with Alkanes TABLE I: Rate Constant Ratios k 3 / k 4for the Gas-Phase Reaction of NO, Radicals with a Series of Alkanes and Ethene at 296 f 2 K
organic n- butane n-pentane n- hexane n-heptane n-octane n-nonane isobutane 2,3-dimethylbutane cyclohexane ethene
relative to n-heotane 0.48 & 0.12 0.59 f 0.12 0.77 f 0.14 1.oo 1.33 f 0.16 1.76 f 0.21 2.95 f 0.13
relative to 2.3-dimethvlbutane
0.24 f 0.05 1.oo 0.33 f 0.06
1.49 f 0.14
a Indicated errors are two least-squares standard deviations of the slopes of plots such as those shown in Figures 1 and 2 combined with the experimental uncertainties arising from the low consumption rates observed.
deviations of the plots of eq I combined with the estimated uncertainties arising from the low consumption rates of certain of these organics. In all cases the intercepts of these plots as derived by least-squares analyses were within two standard deviations of zero. As noted above, several experiments were also carried out to determine the increased rates of decay of N2O5 in the presence of known concentrations of n-heptane and 2,3-dimethylbutane. In all cases the N2O5 decays were observed to be exponential within the experimental error limits, as shown by the linearity of plots of In ([N2O5],/[N2O5],) against time ( t - to). For this technique, the background N 2 0 5decay rate k5 (Le., that in the absence of an added organic) was monitored by FT-IR spectroscopy for -45 min; then, a known concentration of the organic was added to the N205-N02-air mixture and the increased N2O5 decay rate determined.l 4 The resulting data were in good accord with eq 11, and using the equilibrium contant K = 1.33 X T/300)0.32e"080/Tcm3 molecule-' (1.87 X lo-" cm3 molecule-' at 298 K) as given by Malko and Troe,17least-squares analyses yielded the rate constants k3 given in Table 11. The uncertainties in the equilibrium constant K due to the measured temperature changes during the experiment (1 K changes the equilibrium constant by 10%") were negligible (52%). The background N205 decay rates k5 were typically s-' at a relative humidity of 55%, observed to be -1.5 X in good agreement with previous experimental data from this 5800-L chamber. I43l8 Discussion
As seen from Figures 1 and 2, the data obtained by using the relative rate techniques are, within the experimental errors, in good accord with eq I. The rate constant ratios k3/k4given in Table I may be placed on an "absolute" basis (but still linearly dependent on the value of the equilibrium constant K used) using the rate constant for ethene determined previ0us1y.l~ The rate constants k3 so derived, which are the first to be reported for the alkanes, are given in Table I1 under the column "from relative rate technique". The rate constants for n-heptane and 2,3-dimethylbutane given in Table I1 as determined from the enhanced rates of decay of N205in N205-N02-alkane-air mixtures are factors of 1.8 and 2.4, respectively, higher than the rate constants obtained from the relative rate technique. This discrepancy could be due to the presence of minor amounts of alkene impurities in these alkanes, since the I C , alkenes react with NO, radicals several orders of magnitude faster than do the alkanes."J2J4 Thus, for example, since 2,3-dimethyl-2-butene (the most likely alkene impurity in 2,3-dimethylbutane) reacts with NO3 radicals with a rate constant then the of -3 X lo-" cm3 molecule-'^^^ at room temperat~re,'~ presence of 2,3-dimethyl-2-butene at only a 0.0005% level would
-
-
TABLE 11: Rate Constants k 3 for the Gas-Phase Reaction of the NO1 Radical with a Series of Alkanes at 296 f 2 Ka iOi7kl, cm3 molecule-' s-I from N205 from relative decay ratesb rate techniquec alkane 2.0 f 1.0 n-butane 2.4 f 1.2 n-pentane 3.2 f 1.5 n-hexane 7.3 f 2.0 4.1 i 1.8 n-heptane 5.5 f 2.5 n-octane 7.2 f 3.3 n-nonane 2.9 1.4 isobutane 29 f 6 12.1 i 5.4 2,3-dimethylbutane 4.0 f 1.9 cyclohexane
*
aAll data are based upon the equilibrium constant as evaluated by Malko and Troe17 and are linearly dependent on this equilibrium constant. Indicated error limits are two least-squares standard deviations of the slopes of the plots of eq 11. CPlacedon an "absolute" basis using the rate constant for ethene of (6.1 f 2.6) X cm3 molecule-' s-'.14 enhance the N 2 0 Sdecay rate by the factor of 2.4 observed in this study. While this impurity level is below the detection limit of our gas chromatographic technique, it should be noted that the enhanced N2O5 decay due to the presence of such an alkene impurity would only be important during the initial stages of the experiments, since the impurity would be rapidly consumed due to reaction with the large excess source of NO, radicals under the present experimental conditions. For example, in a typical experiment with initial 2,3-dimethylbutane, NO2, and N 2 0 5 5 X lo',, and 2 X lo1, molecules ~ m - ~ , concentrations of 5 X respectively, a 0.0005% level of 2,3-dimethyl-2-butene impurity amounts to only 2.5 X lo9 molecules cm-, and would have a half life of 5 1 s relative to reaction with NO, radicals14 (present at concentrations of -2 X 1O'O molecules cm-, 17). Since the N 2 0 5 decays were measured for periods of the order of 50-90 min, such a low alkene impurity level would only have an effect during the initial stages of these experiments, and larger levels of less reactive impurities would be expected to cause nonexponential decays, which were not observed. A more probable cause for the higher rate constants derived from the N2O5 decay measurements arises from consumption of N2O5 caused by reaction of NO, radicals with the products of the initial reaction of NO, radicals with the alkanes. Although these products have not been determined to date, potential products such as the aldehydes react with NO, with a rate constant of 1.4 X cm3 molecule-' s - I . " , ' ~ Thus, if such compounds are formed, enhanced N 2 0 5 decay rates would be expected. It is apparent that for certain organics or classes of organics the N2O5 decay rate technique is susceptible to secondary reactions, and the rate constants obtained for these organics may be upper limits to those for the initial reaction. However, the relative rate technique is not affected by this problem, and hence the rate constant data obtained with this technique are to be preferred for the alkanes and, as discussed previo~sly,'~ also for the aromatic hydrocarbons. The present data, which are the first to be reported for the reaction of NO, radicals with alkanes, show that, in a manner analogous to the corresponding reactions of O(,P) atomsz2and O H radicals,23 the rate constants increase monotonically along the n-alkane series. These reactions must proceed via H atom abstraction
-
NO,
+ RH
HNO,
-+
+R
with this pathway being exothermic by 3.2,6.7, and 9.2 kcal mol-' for primary, secondary, and tertiary C-H bonds, re~pectively.~~ If one makes the reasonable assumption that H atom abstraction (22) Herron, J. T.; Huie, R. E. J . Phys. Chem. Ref. Data 1973, 2, 467. (23) Atkinson, R.; Carter, W. P.L.; Aschmann, S. M.: Winer, A. M.; Pitts, J. N., Jr. Int. J . Chem. Kiner., in press.
2364
J . Phys. Chem. 1984, 88, 2364-2368
from primary -CH3 groups is slow compared to abstraction from secondary -CH,- groups, the rate constant data for the n-alkanes and cyclohexane allow a group rate constant of (9 2) X cm3 molecule-' s-' to be derived for the reaction of NO, radicals with -CH2- groups. The present data base is not sufficiently precise to allow any further differentiation into -CH2- groups bonded to -CH2- or -CH, groups, as carried out recently by Atkinson et al.23for the reactions of OH radicals with a series of alkanes. The rate constants for isobutane and 2,3-dimethylbutane allow >CH- group rate constants of k[CH(CH,),] = 2.9 X cm3 molecule-' s-' and k[CH(CH)(CH,),] = 6.0 X cm3 molecule-' s-l to be derived at room temperature (where k[CH(CH,),] and K[CH(CH)(CH,),] refer to the rate constants for NO3 radicals reacting with a >CH- group bonded to three -CH, groups and one >CH- and two -CH3 groups, respectively). That these >CH- group rate constants are higher than those for -CH2groups is expected by analogy with O H radical reactionsz3 and with the fact that H atom abstraction from a >CH- group is 2.5 kcal mol-' more exothermic than is H atom abstraction from a -CH2The rate constants obtained in this work show that for nighttime NO3 radical concentrations of -2 X lo9 molecules cm-3 (Le., 100 ppt, a value we have observed in urban and desert areas of C a l i f ~ r n i a ~ these ~ ~ ~ ' nighttime ~) reactions are an order of magnitude or more less effective as total loss processes for the alkanes compared to daytime reaction with the O H radical (assuming an average daytime O H radical concentration of 1 X lo6 molecules ~ m - ~ However, ). these nighttime reactions of alkanes with NO3 do lead directly to loss of NO, and concurrent formation of nitric acid, a key species involved in acid deposition.
*
-
-
(24) Benson, S. W. "Thermochemical Kinetics", 2nd ed.; Wiley: New York, 1976.
Thus, for example, in polluted urban atmospheres nighttime NO3 radical and NO2 concentrations are -2 X lo9 and (5-7) X 1O'O molecules cm-, (- 100 ppt and 2-3 ppb) re~pectively,~ leading" to an N20Sconcentration of -2 X lo9 molecules cm-, (-80 ppt). For an alkane concentration of 100 ppb, this then results in a HNO, formation rate of -0.04 ppb h-l from the reaction of NO3 radicals with n-alkanes. This rate can be compared to a HNO, formation rate of -0.3 ppb h-I from the hydrolysis of NzOs with H 2 0 at -50% relative humidity by using the upper limit value for the homogeneous gas-phase rate constant of the reaction of N205 with water vapor recently determined in these laboratories.18 For air masses typical of semiarid and desert area^,^,'^ NO, radical and NO, concentrations are 5 2 X lo8 and -2 X 1O'O molecules cmT3,respectively. For this situation, if one assumes an alkane concentration of -50 ppb, the H N 0 3 formation rates from the reaction of NO3radicals with alkanes and from the N2O5 hydrolysis reaction are both lower by a factor of -20-30 than those for the polluted atmosphere scenerio described above. Thus, for both polluted and relatively unpolluted atmospheres the rate of formation of HNO, from the reaction of NO, radicals with alkanes can be a significant fraction of the total nighttime HNO, formation rate. Clearly, the reactions of NO3 radicals with alkanes should be included in chemical models of regional air pollution for the assessment of the impacts of acid deposition.
-
Acknowledgment. We gratefully acknowledge the financial support of the U S . Environmental Protection Agency Grant No. R807739-02 and wish to thank Sara M. Aschmann for carrying out the gas chromatographic analyses and W. D. Long for assistance in conducting the chamber experiments. Although the research described in this article has been funded by the Environmental Protection Agency, it has not been subjected to Agency review and therefore does not necessarily reflect the views of the Agency, and no official endorsement should be inferred.
A Photoelectron Spectroscopic Study of the Ground States of CH2F+ and CD,F+ Lester Andrew%*John M. Dyke, Neville Jonathan, Noureddine Keddar, Alan Morris, and Abed Ridha Department of Chemistry, The University, Southampton. SO9 5NH, U.K. (Received: July 14, 1983)
The fluorine atom/methyl fluoride reaction has been studied by photoelectron spectroscopy. A new product band with vibrational components at 9.04 0.01 eV adiabatic and 9.22 0.01 eV vertical ionization energies is assigned to the CH2Ffree radical. The v'= 0-1 vibronic separation measured as 1450 f 30 cm-' is due to the C-F stretching fundamental of the ground state of CH,Ff; this vibronic interval was 1530 f 30 cm-' for CD2F+. The positive deuterium shift for CH2F+is due to interaction with the H-C-H bending mode, which shifts below the C-F stretching mode on deuteration. The substantial increase in the C-F stretching modes for CH2F+and CD2F+,as compared to 1163- and 1193-cm-' values for the CH2Fand CD2F free radicals in solid argon, respectively, is due to increased net C-F bonding in the cations.
*
Introduction The application of photoelectron spectroscopy to the study of transient species has increased in recent years as methods for producing free radicals, high-temperature molecules, and hydrogen-bonded complexes have been Recently, the unstable diatomic fluorides BF, CF, NF, and O F have been prepared for photoelectron spectroscopicstudies of the ground-state cations using evaporative methods and fluorine atom reaction^.^^ The larger fluoromethyl cations have been studied to date only by mass spectroscopic and infrared absorption matrix methods.'&12 In particular, the most intense infrared absorptions of CF3+, *To whom correspondence should be addressed at the Chemistry Department, University of Virginia, Charlottesville, VA 22901.
0022-3654/84/2088-2364$01.50/0
*
CF2C1+, and CHFCl+ have been recorded in solid argon matrice^.'^-'^ These cations exhibited unusually high C-F stretching (1) Dyke, J. M.; Dunlavey, S. J.; Jonathan, N.; Morris, A. Mol. Phys. 1980. 39.- ,1 1 21. -~- - , -
(2) Dyke, J. M.; Jonathan, N.; Morris, A.; Winter, M. J. Mol. Phys. 1981, 44. . ., 1059. (3) Andrews, L.; Dyke, J. M.; Jonathan, N.; Keddar, N.; Morris, A.; Ridha, A. Chem. Phys. Lett. 1983, 97, 89. (4) Colbourn, E. A.; Dyke, J. M.; Lee, E.; Morris, A,; Trickle, I. R. Mol. Phys. 1978, 35, 873. ( 5 ) Carnovale, F.; Livett, M. K.; Peel, J. B. J . Am. Chem. SOC.1982, 104, 5334. (6) Dyke, J. M.; Kirby, C.; Morris, A. J . Chem. SOC.,Faraday Trans. 2 1983, 79, 483. (7) Dyke, J. M.; Lewis, A. E.; Morris, A. J . Chem. Phys., 1984, 80, 1382.
0 1984 American Chemical Society
