J. Phys. Chem. 1987, 91, 1881-1883 with anthracene itself, probably for stereoelectronic reasons. Thus, aromatic cations which are stabilized by extensive charge delocalization or by donor substituents will suffer diminished charge-transfer interactions with neutral aromatics.l* Indeed alkylbenzenes, naphthalene, and pyrene all form dimer cations more readily than a n t h r a ~ e n e . ~ ~ . ~ ~ Finally, we wish to emphasize that the difference between An+ and An2+ bears on the more general questions relating to concentration dependences on reactivity in which the deviations from the dilute solution approximation are customarily dealt with by activity coefficients. For example, reactions such as dimerizations and cycloadditions can be accelerated by association at high concentrations. On the other hand, the ionic additions addressed here (see eq 2) lead to decreased reactivity at higher concentrations. In order to compensate properly for such diverse behavior,
1881
a consideration of activity coefficients which are reaction specific as well as species specific is required. We feel that a knowledge of the microscopic processes of the kind described here will simplify the interpretation of associative interactions. These are especially important in determining ionic reactivity, and they will hopefully allow valid kinetic and thermodynamic information to be derived.39437~18320 We are currently engaged in studies directed toward this goal. Acknowledgment. We thank S. J. Atherton and M. A. J. Rodgers of the Center for Fast Kinetics Research (under support from N I H Grant RR00886 and the University of Texas, Austin), the Sohio Fund and Cleveland State University for financial support (J.M.M.), and grants from the National Science Foundation and Robert A. Welch Foundation. Registry No. An, 120-12-7; Ant, 34512-28-2; Anz+, 75209-13-1; T N 0 2 , 509- 14-8; [An+,T-], 106 139-49-5; [An,+,T-], 10616 1-09-5.
(18) See: Foster, R. Organic Charge-Transfer Complexes; Academic: New York, 1969. Zoos, Z. G.; Klein, D. J. Mol. Assoc. 1975, 1 , 2. Moet-Ner, M.; El-Shall, M. S. J . Am. Chem. SOC.1986, 208, 4386. (19) Masnovi, J. M., unpublished results.
Rate Constant for the Reaction I
+ F,
-
(20) Friend, K. E.; Ohnesorge, W. E. J . Org. Chem. 1963, 28, 2435. Parker, V. D. Ace. Chem. Res. 1984, 17, 243.
IF
+F
H. V. Lilenfeld* and G. R. Bradburn McDonnell Douglas Research Laboratories, St. Louis, Missouri 63166 (Received: August 20, 1986; In Final Form: December 2, 1986)
-
The rate constant for the reaction I + F2 IF + F was measured at room temperature in a discharge-flow apparatus. The value obtained was (1.9 & 0.4) X cm3 molecule-’ s-’. An upper limit for the rate constant for the reaction between CF31+ F2 was estimated from the experimental data to be 1 X IO-’’ cm3 molecule-’ s-’.
Introduction The reactions of iodine atoms with halogen molecules are of considerable interest, partially because of their involvement in chemical mechanisms which produce chemiluminescence. For example, it has been suggested that the reaction I
+ Fz- ki
IF
+F
is involved in the mechanism of production of electronically excited I F (IF*) when 0 2 ( I A ) is added to a mixture containing IF.’ In addition, mechanisms which have been advanced to explain the chemiluminescence observed when I2 and F2 react in flowing systems have included reaction 1 These mechanisms have been partially confirmed by the detection of atomic iodine in a flowing mixture of I2 and-F2.4 Understanding of the kinetics of I F formation in these systems requires a knowledge of the rates of the reactions involved in the mechanism. Attempts to accurately measure the rate of reaction 1 by use of flash photolysis in slow-flow systems, however, have been thwarted by the presence of a dark reaction between F2 and the perfluoroalkyl halide precursor used in these experiment^.^ Our experiments, as described below, performed in a fast-flow reactor under pseudofirst-order conditions, are the first accurate measurements of the rate of reaction 1. ,233
Experiment
The experimental apparatus used in these experiments has been described previously in detaiL6 A brief description of the flow (1) Whitefield, P. D. J . Photochem. 1984, 25, 465. (2) Kahler, C. C.; Lee, Y. T. J. Chem. Phys. 1980, 73, 5122.
(3) Coggiola, M. J.; Valentini, J. J.; Lee, Y. T. Int. J . Chem. Kinet. 1978, 8, 605. (4) Lilenfeld, H. V.;Bradburn, G. R. Chem. Phys. Lett. 1986, 131, 276. (5) Berman, M. R.; Whitefield, P. D. J . Chem. Phys. 1986, 84, 4281.
0022-3654/87/2091-1881$01.50/0
system and of the modifications to the above apparatus for the experiments described herein follows. Flow System. The flow system consists of a 2.5-cm-diameter flow reactor with electron paramagnetic resonance (EPR) and optical emission and mass spectrometric detection systems as shown in Figure 1. The system is pumped by a 140-L/s pump and flow velocities of up to 20 m/s can be obtained. A reactive chemical species (usually a radical) in an inert diluent is introduced at point A (Figure 1) in the flow. A second species, whose concentration is in excess of the radical, is injected through a movable injector at any point along the flow axis. The radical species, and possibly the products of the reaction, are detected as a function of the injector-to-detector distance. The pressure in the flow reactor is measured with a capacitance manometer and the flow rates of the reactants and diluents are measured with calibrated flow gauges. Atomic Iodine Source. The major modification to the experimental apparatus for these experiments was the development and incorporation of an atomic iodine source. The atomic iodine was formed by passing a mixture containing a few percent CF31 in Ar through a microwave discharge. The effluent of the discharge was studied with EPR, optical spectroscopy, and mass spectroscopy. The discharge power was adjusted such that only a small fraction of the CF31was dissociated. Under our experimental conditions, atomic iodine was detected in the discharge effluent with EPR, but atomic fluorine was not. The visible emissions from the microwave discharge extinguished themselves within a few centimeters downstream of the discharge cavity. No emissions from either the I(2P,12) I(2P312)infrared transition at 1.315 pn or from the known IF(A X) or IF(B X) visible transitions were observed with a 0.25-m monochromator/detector
--
-
(6) Lilenfeld, H. V.; Whitefield, P. D.; Bradburn, G. R. J . Phys. Chem. 1984, 88, 6158.
0 1987 American Chemical Society
1882 The Journal of Physical Chemistry, Vol. 91. No. 7, 1987
-
MW smometa
Lilenfeld and Bradburn TABLE I: Determination of the Rate Constant for the Reaction I F2 IF + F 10-15[ F2], d In ([I]/[I]o)/ velocity, molecule/cm’ [F,]/[I], dt, s-’ mls
\
A
1.58 0.70 0.91 1.60
Microwave cavity cavity
1.09 1.37 Movable
Light pipe to 0.25-m monochromator/detcctor
injector
17.5 11.8 18.2 8.9 9.4 18.8
24.3 8.2 12.2
+
11.0
14.9 14.9
26.6
9.3
17.5
13.5
22.4
10.8
x
ml
B
Figure 1. Fast-flow apparatus.
system positioned approximately 10 cm downstream of the discharge region.
Results Previous investigations have shown the presence of complicating reactions when mixtures of fluorine with perfluoroalkyl halides were ~ t u d i e d . ~These , ~ investigations, however, were performed under slow-flow conditions and indicated that the complicating reactions became less important at higher flow velocities. These complicating reactions were therefore studied in our fast-flow apparatus to ascertain whether they could influence our experiments. The results of these experiments, as well as the determination of k , , are discussed below. A. CFJ F2. To determine the rate of the CF,I F2reaction, F2 in He carrier gas was added to a flow of CF,I in Ar with the movable injector. CFJ was monitored mass spectrometrically via its parent peak ( m / e 196, electron energy = 20 V) as a function of injector-to-detector distance. Under these fast-flow conditions no reaction was observed. These results indicate that an upper limit of 1 X cm3 molecule-’ s-I exists for the rate constant for this reaction. B . I + F,. The rate of reaction 1 was measured with the atomic-iodine source described above. A 10% mixture of F2 diluted in H e was added to the effluent from the atomic-iodine source through the movable injector. Atomic iodine was monitored as a function of injector-to-detector distance (proportional to reaction time) by use of EPR. In addition, CF31 and F2 concentrations were monitored with mass spectroscopy. These experiments were performed under pseudo-first-order conditions (excess F2). The measured CFJ and F, concentrations did not change appreciably during the course of these experiments. The atomic-iodine concentration, however, decreased as a function of reaction time. No atomic fluorine was observed during the course of the reaction, presumably because the atomic fluorine formed by reaction 1 is consumed by the rapid subsequent reaction
+
+
CFiI
+F
k2 +
CF3
+ IF
(2)
The presence of excess CF31 in the system therefore rapidly converts atomic fluorine to iodine monofluoride. Thus, this process acts as a chemical “getter” for atomic fluorine which otherwise might react with small levels of I2 impurities (formed at the walls either by I-atom recombination or by I F decomposition) to reproduce atomic iodine by the reaction F
+ 1 2 - ki
IF
+I
(3)
The atomic-iodine concentration was monitored as a function of injector-to-detector distance for various flow velocities and fluorine concentrations, where [F2] >> [I]. The procedure used to derive the rate constant for reaction 1 from the data has been described in detail by Westenberg and De Haas.* It involved measurement of [I] at various detector positions with a measured flow of the F,/He mixture. Subse(7) Burrows, M. D. J. Chem. Phys. 1984, 81, 3546. (8) Westenberg, A. A.; De Haas, N. J. Chem. Phys. 1967, 46, 490.
0.0 2.0 4.0 6.0 8.0 10.0 12.0 14.0 16.0 18.0 20.0
[F,] (molecules/cm3)
*1d4
Figure 2. The rate of reaction of I as a function of [FJ.
quently, the F2/He flow was either injected at a point in the flow beyond the EPR detector or replaced by an equivalent flow of He. The logarithmic rate of reaction of iodine, defined as the derivative of In ([I]/[I]o) with respect to reaction time, where [I] and [I],, are the concentrations of iodine with and without F2 present in the flow, was calculated from the data. As shown previously,6 this procedure corrects for any atomic iodine production or reaction by species present in the effluent of the discharge (flow A in Figure 1). The only reactions measured in this procedure are those which involve F2 or reaction or consumption of I by the products of reaction 1. The logarithmic rate of reaction of iodine was determined in several experiments with differing flow velocities and fluorine concentrations. The results of these experiments are shown in Table I. A plot of the logarithmic rate of reaction of iodine as a function of fluorine concentration is shown in Figure 2. The line shown in Figure 2 is a nonweighted, least-squares fit of the data; the line has a slope of (1.94 f 0.37) X cm3 molecule-’ s-l, and an intercept of -4.9 f 4.6 s-’ (the reported errors are 95% confidence limits). The presence of a negative intercept in our results suggests the possibility of the presence of a secondary, atomic-iodine-producing process in the system. As explained above, the experimental procedure used in our experiments automatically corrects for production or consumption of iodine by any reactants present in the effluent of the CFJ discharge. Therefore, only reactions involving F2 or the products of reaction 1 should influence our experimental results. It has been suggested by wevious workers5 that the reaction CF3 + I F
-
CF4
+I
(4)
might account for atomic-iodine formation in CF,I/F2 systems. Therefore, another experiment was performed to determine if reaction 4 could influence our results. C. CF, IF. Atomic fluorine was formed by passing a mixture of F2 in He through the microwave discharge. CF31 in Ar was added to the flow with the movable injector. The conditions of this experiment (i.e., the flow rates and velocity) were similar to the conditions of the experiments in part B. The reaction between CF31and F is known to proceed through the reaction path
+
CF31 + F
-
CF,
+ IF
(5)
Given the similarity of the conditions between this experiment and those in part B, if reaction 4 were fast enough to influence
J . Phys. Chem. 1987, 91, 1883-1887
+
the experiments in part B, we would expect the CF3 IF formed in reaction 5 to form detectable amounts of atomic iodine in the present experiment; however, no atomic iodine was detected. Conclusions
The previous attempts to measure the rate of reaction 1 by using laser photolysis to form atomic iodine in slow-flow systems were unsuccessful because of the presence of a dark reaction between the precursor, CFJ, and F,. This reaction was, however, too slow to interfere with the kinetics in our experiment where atomic iodine was formed by discharging a mixture of CF31in Ar in a fast-flow apparatus. We estimate the upper limit for the dark reaction to be 1 X IO-'' cm3 molecule-' s-l. The rate constant for reaction 1 was determined to be 1.9 X cm3 molecule-' s-I from the slope of a plot of the pseudofirst-order rate constant for reaction 1 as a function of [F2I0. The negative nonzero intercept of the plot suggests that an iodineproduction reaction may be present in the system; however, no experimental evidence of atomic-iodine production was seen when atomic fluorine and CF31were reacted. An alternative explanation of the negative intercept might be a change in the rate of wall recombination of I atoms caused by adding F2 to the flow. We estimate the recombination rate for 1, on the surface of the reactor to be of the order of 5 s-I. Any significant reduction of this rate when F2is added (i.e., a poisoning of the walls by F,) could cause the negative intercept observed in our experiments. Our value for k l is in good agreement with an estimate, k l F= 1 X cm3 molecule-' s-I, for this rate constant by Berman and Whitefield5 from their laser photolysis experiments. The reaction pair I(2Pl/2)
+ Fz
-
IF
+F
(6)
and reaction 1 offer an interesting comparison for theoretical analysis because both reactions involve an exothermic pathway
1883
for fluorine displacement by atomic iodine. The recently measured cm3 rate for the deactivation of I(2Pl/2)by F2s,9( k = 5 X molecule-' s-'), when compared with our measured rate for (I), suggests that the rate constants for these reactions may be of the same order of magnitude. Calculations have been performed by using potential energy surfaces for the reaction pair F(2Pl,z) + H,
+ H2
F('P312)
-
products
(7)
+H
(8)
HF
which show that reaction 8 proceeds at a much faster rate than reaction 7. In addition, experiments indicate" that the rate constant for the reaction Br(2P3,2)
+ HI
-
HBr
+I
(9)
may be greater than the rate constant for the deactivation reaction
+ HI
Br(ZPl,2)
products12
(10)
Calculations on the potential energy surface for reactions 1 and 6 and a measurement of the rate constant for reaction 6 would afford another chance for comparison of experiment with theory. Acknowledgment. This research was conducted under the McDonnell Douglas Independent Research and Development Program. We thank Drs. M. R. Berman and P. D. Whitefield for helpful comments during the writing of this manuscript. Registry No. I, 14362-44-8;F2,7782-41-4;CF31, 23 14-97-8. (9) An upper limit for this reaction was also recently obtained in Chowdhury, M. A,; Pritt, Jr., A. T.; Patel, D.; Benard, D. J. J. Chem. Phys. 1986, 84, 6687. (IO) Tully, J. C. J . Chem. Phys. 1974, 60, 3042. (11) Bergmann, K.; Leone, S. R.; Moore, C. B. J . Chem. Phys. 1975,63, 4161. (12) For an analysis of these results see, Houston, P. L. Chem. Phys. Lett. 1977, 47, 137.
Maximally Inhibited Pyrolysis Kinetics of Bromo Esters in the Gas Phase. The Ion Pair Mechanism Gabriel Chuchani* and Rosa M. Domhguez Centro de Quimica, Instituto Venezolano de Iavestigaciones CientNcas, Apartado 21 827, Caracus 1020-A, Venezuela (Received: April 28, 1986) The kinetics of the pyrolyses of two bromo esters have been investigated in a static system over the temperature range 379-419 OC and pressure range 51-1 10 Torr. The reactions in seasoned vessels, and under maximum inhibition with the radical chain suppressor propene, are homogeneous, unimolecular, and follow a first-order rate law. The Arrhenius equations for these elimination processes were found to be as follows: for methyl 4-bromobutyrate, log k, (s-I) = (13.38 A 0.59) - (216.7 i 7.6) kJ mol-' (2.303R7')-I; for methyl 5-bromovalerate, log k l (s-l) = (13.84 0.17) - (228.9 2.3) kJ mol-' (2.303RT)-'. The anchimeric assistance of the carbomethoxy substituent is determinant in the elimination process, where dehydrobromination product and lactone formation result from an intimate ion pair type of mechanism. The partial rates toward each of the primary products have been determined, reported, and discussed. The present work provides further support of the intimate ion pair mechanism in the gas-phase pyrolysis of special types of organic halides.
*
*
Introduction
The pyrolysis of ethyl 4-bromobutyrate in the gas phase' led to several parallel elimination paths as represented in eq 1. The COOCH2CH3substituent was believed to assist anchimerically the elimination process of paths 1 and 2, where dehydrobromination and lactone formation arise from an intimate ion pair mechanism. This study corroborated the work on the gasphase elimination of methyl esters of w-chlorocarboxylic acids.,
BrCH2CH2CH2COOCH2CH3
\ I3
BrCHzCH2CHzCOOH
CH2-Cp2 C=O CH2-0'
I
t CHSCH2Br
CH2=CHCH2COOCH2CH3
+
t HBr
(1)
CH2=CH2
J4
CH2-CH2
(1) Chuchani, G.; Domhguez, R.M.; Martin, I. React. Kine?. C a r d . Lett. 1983, 30, 77. Chuchani, G.; Dominguez, R. M. Int. . I Chem. . Kine?. 1983,
15, 795. (2) Chuchani, G.; Domhguez, R. M.; Rotinov, A. Int. J. Chem. Kine?. 1986, 18, 203. 1982, 24, 381.
0022-3654/87/2091-1883$01.50/0
I
'C=O
t HBr
CH2-0'
Here also, rate enhancement and formation of cyclic products suggested an ion pair mechanism through participation of the 0 1987 American Chemical Society
