T H E
J O U R N A L
O F
PHYSICAL CHEMISTRY Registered in U.S. Patent Office 0 Copyright, 1975, by the American Chemical Society
VOLUME 79, NUMBER 22
A Reexamination of the Reaction H
OCTOBER 23, 1975
+ HCI S H2 + CI
John E. Spencer and G. P. Glass. Department of Chemistry, Rice University, Houston, Texas 7700 1 (Received May 9, 1975) Publication costs assisted by the Petroleum Research Fund
The reaction between atomic hydrogen and HC1 was studied at 295 K in a fast discharge flow system by monitoring, at various reaction times, EPR spectra of H(2S) and Cl(2P3/2).The experimental results could be interpreted as showing that the ratio of the forward and reverse rate constants for the system H + HCI H2 C1 (1) was equal to the equilibrium constant for this system. However, the results were complicated by a sequence of reactions initiated by atomic chlorine formed in reaction This species reacted with HC1 a t the vessel walls, forming a product, probably HC12, that reacted further to consume a second hydrogen atom. The effect of these secondary reactions was neglected in an earlier paper that reported a ratio for the rate constants, kJk-1, greater than the equilibrium constant by a factor of 2 to 3. No evidence was found for the wall reaction between H and C1 that had been postulated as occurring in this system.
-
+
&.
Introduction In 1968, Westenberg and deHaasl investigated the system H
+ HCl F: H2 + C1
(1)
using a discharge flow apparatus attached to an EPR spectrometer. They measured rate constants for both the forward and reverse reactions at temperatures between 251 and 456 K, and found kJk-1 to be always greater than the equilibrium constant by a factor of 2 to 3. This result is not incompatible with the principle of microscopic reversibility because the measurements were made under conditions far from equilibrium. However, it has generated considerable interest, because it brings into question the traditional practice of estimating rate constants from measurements taken on their reverse reactions. In the last 6 years, several attempts have been made to explain the inequality between k&-1 and the equilibrium constant in terms that are unique to this system. In their original paper, Westenberg and deHaas suggested that their measured value of k-1 was low because of a substantial back reaction from rotationally hot HCl, formed by reaction (-1) in a nonequilibrium distribution. An alternative explanation, based upon a consideration of the role of C1(2P1/2)in the reverse reaction, has been offered by Snider,2 and independently by C l ~ n e These .~ authors examined various mechanisms that involved reactions of both
Cl(2P3/2) and C1(2P1/2)with HQ,together with other reactions interconverting the two spin-orbit states. They were able to reproduce Westenberg's results by assuming that the rate of reaction of C1(2P1/2)with H2 was greater than either the rate of reaction of Cl(2P3/2)with H2, or the rate of interconversion between C1(2P1/2) and C1(2P3/2). Under these conditions, the concentration of C1(2P1/2)is reduced below its equilibrium value, and the rate of the overall reaction with H2 is subsequently decreased. These explanations, which are based on a nonequilibrium distribution of states, have been discussed in detail by Galante and G i ~ l a s o nThey . ~ have concluded that they are incompatible with much of the available experimental and theoretical evidence on the system. For example, if Snider's proposal were correct, the rate of the reverse reaction, (-l), would be increased by the presence of a molecule, such as CC14, that efficiently interconverts C1(2P3/2)and C1(2P1/2). However, a recent measurement5 of k-1, made at 300 K and in the presence of a large excess of C&, yielded a value (1.4X cm3 molecule-l sec-l) that is in substantial agreement with those obtained earlier by Westenberg and deHaas' (1.3 X cm3 molecule-' sec-I), and by Benson et a1.6 (1.2 X cm3 molecule-1 sec-l). This agreement between the various measured values of k-1 led Galante and Gislason4 to critically reexamine measurements of the forward reaction between atomic hydrogen and HC1. They noted that Westenberg and deHaas failed to detect the presence of atomic chlorine,' and sug2329
2330
John E. Spencer and G. P. Glass
gested a possible reason for this; namely, the occurrence of a fast wall reaction between atomic chlorine and atomic hydrogen.
--
wall
H+C1
HC1
(2)
They argued that, if this reaction were sufficiently fast, a second hydrogen atom would be consumed by the product of reaction 1, and the value of kl, estimated from the measured hydrogen decay rate, would be high by a factor of 2. The initial purpose of the work described in this paper was to test the hypothesis of Galante and Gislason by measuring directly the rate of reaction 2 in a system containing no HC1. Unfortunately, no evidence for the occurrence of this reaction was obtained, and thus a broader study of the reaction of atomic hydrogen with HCl was undertaken. This study was performed using a discharge flow system equipped for EPR detection of free radical species. In order to simplify the study, reaction 1 was usually investigated using wall coatings that prevented measurable recombination of both atomic hydrogen and atomic chlorine. Experimental Section The discharge flow apparatus, the EPR spectrometer, and the operating techniques used in this study have been described in detail p r e v i o ~ s l y .The ~ ~ ~20-mm i.d. quartz flow tube was operated a t pressures of 0.50 to 1.50 Torr, and at linear flow speeds of 1500 to 2200 cm/sec. Atomic hydrogen and atomic chlorine were generated 50 cm upstream of the EPR cavity by microwave discharge of dilute (0.1-0.5%) mixtures of H2 or cl2 in argon. In some cases atomic chlorine was generated from atomic hydrogen by the fast reaction H
+ C12
--*
HC1+ C1
(3)
A movable double inlet probe? whose position could be varied between 2 and 25 cm from the point of EPR detection, allowed the simultaneous addition of two different gases to the flowtube, one slightly upstream of the other. Flow rates of HC1 (Matheson 99%), Cl2 (Matheson 99.5%), and H2 were estimated by measuring the pressure drop in bulbs of known volume. Flow rates of argon were measured using a calibrated rotameter. EPR signals of H(2S) and of C1(2P3/2,mI = -%, mJ = -$ lh) were recorded with a Varian E12 spectrometer. C1(2P1/2) was not observed. Absolute concentrations were principally determined by double integration of the EPR spectra, but, on occasion, were determined by the method of first moments,g or from the measured peak height and peak widths.10 Known pressures of 0 2 were used throughout as concentration calibration standards. Transition probabilities for all of the observed species have been rep ~ r t e dAll . ~ measurements were made at 295 f 1 K. In order to prevent wall recombination of atoms, the flowtube was treated with various wall poisons. Among those used were (a) a halocarbon wax,ll (b) a halocarbon wax fluorinated as described in ref 8, and (c) orthoboric acid. In some experiments the tube was cleaned and washed with a 5% solution of HF. The rate of chlorine atom recombination was measured on all of these surfaces. Computer simulation of the reaction system was performed throughout this work. Various model reaction mechanisms were tested by comparing the results obtained by integration of the pertinent rate equations with those obtained from experiment. Numerical integration was per-
-
The Journal of Physical Chemistry, Vol. 79, No. 22. 1975
formed using Euler's method on an IBM Model 370/155 computer. Results and Discussion Wall Reactions of Atomic Chlorine. Wall reactions of atomic chlorine were studied on each of the surfaces described in the Experimental Section. However, in flowtubes coated with either halocarbon wax or fluorinated halocarbon wax, wall decay was too slow to be measurable. Also atom recombination was totally inhibited for a period of several hours in tubes freshly washed with a 5% solution of HF. Therefore, all of the rate measurements described below were made in boric acid coated flowtubes similar to those used by Westenberg and deHaas.1 Chlorine atoms were generated by adding varying amounts of Cl2 to atomic hydrogen. In some experiments H
+ Cl2
HCl
-+
+ C1
(3)
sufficient C12 was added to decrease the hydrogen atom concentration to one-half of its original value; in other experiments a small excess of Clp was used. Atom recombination was monitored by EPR, well downstream of the C12 inlet, at positions corresponding to reaction times of 4-14 msec. In these studies the initial concentration of atomic hydrogen was adjusted to -5 X 1013molecule ~ m - so ~ ,that the conversion of H to C1 by reaction 3 was 99% complete within the first 2-4 msec.I2 This time span was also sufficient to allow any vibrationally excited HCl, formed in reaction 3, to be totally deactivated by collision with atomic hydrogen or atomic ch10rine.l~The reaction of H with ground state HCl is too slow1 to allow more than 2% of any remaining atomic hydrogen to be consumed during the 14 msec during which atom recombination was monitored. The results of two measurements of the chlorine atom decay rate are shown in Figure 1. The points marked (a) were obtained by adding sufficient Cl2 to remove all trace of atomic hydrogen. Those marked (b) were obtained by lowering the Cl2 flowrate until one-half of the original atomic hydrogen remained. First-order decay of atomic chlorine was observed in all experiments. The rate constant for chlorine atom recombination, defined as [(Cl)-l d(C1)/ dt], was measured as 177 sec-' in the absence of atomic hydrogen, and as 93 sec-' in the presence of an equal amount of atomic hydrogen. In this latter case, decay of atomic hydrogen also occurred a t a rate equal to that for the decay of atomic chlorine. These results are well explained by the following mechanism:
c1 H
-
wall
+ Cl2
%C12 HCl
--+
+ C1
(4) (3)
In this sequence, reaction 4 is assumed to be slow and rate limiting, and this allows reaction 3 to regenerate one-half of the originally recombined atomic chlorine, and thus reduce the apparent rate of wall loss by a factor of 2. The sequence predicts an equal loss rate for H and C1. The results are not consistent with C1 being consumed mainly by the wall reaction suggested by Galante and Gisla~on.~
H
-
+ C1 wall HCl
If this reaction were dominant, the rate of recombination of
2331
Reaction between Atomic Hydrogen and HCI I
I
I
TABLE I: Stoichiometry and Rate of Reaction of Atomic Hydrogen with HCla (Cl)/[(H), - (H)] (HC1)d (HCl),
At 5 msec
At 10 msec
(H),
(H),
k,,
(a) In 7.28 5.96 3.52 1.81
an HF Washed Tube 11 7.48 16 0.33 0.26 8.20 25 0.16 0.12 9.71 46 0.19 0.11 9.53 Av 8.7 * 0.9
~~~~~
82.5 92.1 87.1 83.5
(b) Flowtube Coated with Fluorinated Halocarbon Wax TIME
(rnsec)
Figure 1. Chlorine atom decay along a boric acid coated flow tube: (a) atomic hydrogen absent; (b) in the presence of an equal initial amount of atomic hydrogen. The points marked 0 refer to atomic hydrogen. atomic chlorine would increase in the presence of atomic hydrogen. The Reaction of H with HC1. A broader study of the reaction of H with HC1 was initiated when it became clear that the hypothesis of Galante and Gislason concerning reaction 2 was incorrect. In order to avoid the complications associated with wall recombination of C1, all studies of this reaction were carried out in flowtubes coated so as to prevent any measurable recombination of atomic chlorine or atomic hydrogen. Frequent checks were carried out as described in the above section to ensure that these flowtubes remained inert throughout the study. In the initial experiments, a large excess of HC1 was added to atomic hydrogen and the subsequent H atom decay and C1 atom growth were followed, using EPR, at reaction times varying from 2 to 15 msec. The decay of atomic hydrogen was found to be approximately first order, as expected, but the amount of C1 produced was found to be always less than the amount of H consumed. Furthermore, the rate of H atom loss, and the ratio of C1 formed to H consumed, was found to be dependent upon the specific wall coating used. Some of the pertinent data from these experiments is shown in Table I. A close examination of this table reveals a correlation between the rate of hydrogen atom decay and the ratio of C1 formed to H consumed. In an HF washed flowtube, the average rate constant [k,,, = -(HCl)-l cm3 molecule-' d(H)/dt] was estimated as 8.7 X sec-l, and the average ratio of C1 formed to H used as 0.20. A rate constant of 7.3 X and a stoichiometric ratio of 0.31, was obtained when using a flowtube coated with fluorinated halocarbon wax, while values of 5.8 X and 0.78 were determined for the rate constant and the stoichiometric ratio, respectively, when using a flowtube freshly coated with normal halocarbon wax. Although there is some scatter in individual experiments, it is clear that the stoichiometric ratio decreases as the rate of hydrogen atom removal increases. An extrapolation to a stoichiometric ratio of unity yields a value for kappof 4.3 x cm3 molecule-l sec-l. These results may be compared with those obtained by previous workers. Clyne and Stedman14 measured the apparent rate constant as 7.3 x 10-14 cm3 molecule-' sec-1,
106.0 95.0 69.2 64.3 135 85.8 105
18.9 15.9 1.02 7.12 14.8 9.26 7.99
5.6 6.0 8.6 8.9 9.1 9.3 13
0.30 0.40 0.34 0.27 0.31 0.34
7.13 8.29 6.75 0.24 7.33 0.30 7.27 0.28 7.13 0.29 7.19 Av 7.3 f 0.5 0.32
Freshly Coated with Wax 5.4 0.70 0.74 6.43 7.3 0.71 0.78 5.93 9.7 0.68 0.78 5.53 12 0.65 0.77 5.64 18 0.87 0.84 4.90 45 0.57 0.68 6.12 Av 5.8 f 0.5 a All concentrations are in units of 1013 molecule/cm3. k RIj,, decm3 fined as -[d In (H)/dt]/(HCl),are recorded in units of 118 126 107 172 128 135
(c) Flowtube 21.7 17.2 11.0 14.1 7.26 2.97
molecule-1 sec-I.
while Westenberg and deHaas' obtained a value of 11 X cm3 molecule-l sec-l. Neither group observed atomic chlorine, although the apparatus used by Westenberg and deHaas should have been sensitive enough to have detected it, if it had been present in significant quantity. The rate constant for hydrogen atom decay may also be compared with the value of 4.3 X cm3 molecule-' sec-', estimated for 121 by setting kJk-1 equal to the equilibrium constant, and assuming a value for 12-1 that is equal to the average of three recent determination^.'^^^^ It is interesting that this value is equal to that obtained by extrapolating our results to a stoichiometric ratio of one. The dependence of kapp on the stoichiometric ratio suggests that chlorine atoms, formed in reaction 1, react to increase the overall rate of hydrogen atom decay, while at the same time, are themselves consumed. A direct test of this hypothesis was made by measuring the effect on the hydrogen atom decay rate of extra added atomic chlorine. In these experiments, atomic chlorine was generated well upstream of the HCl inlet by titrating approximately onehalf of the initial H atom concentration with Clz. A large excess of HC1 was added, and the remaining atomic hydrogen was monitored by EPR at various positions downstream from the HCl inlet. A marked acceleration in the rate of hydrogen atom decay always occurred when C1 was added to the system. This is illustrated in Figure 2. Rate constants, estimated from the approximately first-order decay of atomic hydrogen, are listed in Table I1 for experiments performed in both the presence and absence of extra atomic chlorine. As can be seen, the rate of disappearance The Journal of Physical Chemistry, Vol. 79, No. 22. 1975
John E. Spencer and G. P. Glass
2332
5
-‘u - 20l u
m ” 5
-
I-
I t 3
TIME
I
9
6
(rnsec)
TABLE 11: Effect of Added Atomic Chlorine on the Rate of Removal of Atomic Hydrogen in the H-HC1 Systema (HIo
(Cl),
Flowtube coating ~~
b
k,, ~
~~
(a) 82.5
7.28 0 HF washed 7.48 3.6 H F washed 28.4 77.1 2.93 (b) 135.0 14.8 0 Fluorinated wax 7.27 7.9 Fluorinated wax 12.8 6.9 135.0 (c) 106.0 18.9 0 Fluorinated wax 7.13 12.0 6.9 Fluorinated wax 11.8 106.0 a Concentrations are in units of 1013 molecule/cms. k,,,,, defined cm3 moleas-[d In (H)/dt]/(HCl),arerecorded in units of cule-1 sec-1.
of atomic hydrogen was approximately doubled by the addition of C1. In an effort to explain these results, all possible reactions between C1 and the other major constituents present in the system were examined in isolation. Molecular hydrogen was found to react slowly with C1, but it produced rather than removed atomic hydrogen. Atomic hydrogen itself was not affected by the presence of C1. However, when a large excess of HC1 was added to atomic chlorine, produced either by reaction 3 or by microwave discharge of dilute C12-Ar mixtures, the chlorine atom concentration decreased as shown in Figure 3. The pattern of behavior shown in Figure 3 was typical of all our experiments in that the concentration of C1 dropped prior to the first observation point, which was situated a t a position corresponding to a reaction time of about 3 msec, but did not change further a t longer reaction times. This behavior is unusual, but is consistent with equilibrium being established between C1 and HCl. C1+ HCl
products
(5)
In order to test this hypothesis, a large number of measurements were made, in a number of different flowtubes, of the decrease in concentration of atomic chlorine brought about by the addition of varying amounts of HCl. In these experiments, chlorine atom concentrations were calculated from measurements made on the EPR line widths and signal heights,10 since the addition of HC1 tended to broaden the EPR signals. The modulation amplitude was reduced The Journal of Physical Chemistry, Vol. 79, No. 22, 1975
I
I ]
IO
15
(rnsec)
TIME
Figure 2. The effect of extra added CI on the decay of atomic hydrogen in the reaction with HCI: (a) no added CI:the apparent first-order cm3 molecule-1 sec-‘; ( b \ p the rate constant, k, is 7.13 x of added CI: k = 11.8 X 10- Cm3 presence of 6.9 X 1 O l 3 molecule-’ sec-’. Tube coated with fluorinated halocarbon wax.
(HCl),
I
5
Flgure 3. The addition of HCI to atomic chlorine in a tube coated with fluorinated halocarbon wax. X represents measurements in the absence of HCI; 0 represents measurements taken after addition of loi5 molecule cm-3 of HCI.
to the point where the signal height was directly proportional to its magnitude. The results of these measurements, together with estimates of the ratio [(Cl)o - (Cl)/(HCl)Cl)], are shown in Table 111. As can be seen, the ratios fall into cm3 moletwo groups. One, clustered around 2 X cule-I, was obtained with HF washed flowtubes, and with flowtubes coated with aged or fluorinated halocarbon wax. The other group, with ratios lower by a factor of 5, was obtained using flowtubes freshly coated with halocarbon wax. It should be noted at this point that the measured ratio is equal to the equilibrium constant for reaction 5 only when the measured atomic chlorine concentration is at its true equilibrium value. At this point, computer simulation of the system was carried out using the following simple model mechanism: H
+ HC1-
C1+ HCl H
+ HClz
H2
+ C1
* HCl2 +
2HC1
(1) (5) (6)
HClz was chosen as the simplest product of reaction 5, and reaction 6 was included in order that extra atomic hydrogen might be removed. Reactions 7 and 8 were omitted because a number of experiments (see Figure 3) C1+ HCl2 HCl2
+
+ HCl2
HC1+ Cl2
-+
2HC1+ Cl2
(7) (8)
showed that atomic chlorine was not continuously removed in the presence of a large excess of HCl, as it would be if the sequence (5) and (7), or (5) and (8) occurred. Reaction -3 is 45 kcal/mol endothermic, and was not considered. C1+ HCl
-
H
+ Cl2
(-3)
The rate equations pertaining to this mechanism were numerically integrated, using an IBM 3701155 computer, for a variety of initial conditions similar to those used in the experiments listed in Tables 1 and 11. The following constraints were placed on the rate constants chosen for the above reactions. (a) k 1 was set equal to 4.3 X lod1* cm3 molecule-l sec-1, the value estimated from the equilibrium constant and the measured rate constant for the back react i ~ n . l ,(b) ~ . ~For the simulation of experiments performed in flowtubes washed with 5% HF, or coated with fluorinated halocarbon wax, the ratio k5lk-5 was set equal to 2.1 X
2333
Reaction between Atomic Hydrogen and HCI TABLE 111: Effect of Adding Large Amounts of HCl to Atomic Chlorinea
TABLE IV: Results of Computer Simulationsa (CW
10'6
[(H)o- (HI] (HCl),
7.73 7.58 7.58 7.44 7.04 1.13 4.33 4.25 4.09 3.28 5.27 5.63 14.3 14.O 13.2
(a) Flowtube Washed 7.35 0.38 6.67 0.91 6.53 1.05 5.71 1.73 4.62 2.42
with HF 32 .O 77.6 84.7 140 310
1.6 1.8 1.9 2.2 1.7
(b) Flowtube Coated with Fluorinated Wax 0.75 0.38 225 2.2 3.87 0.46 64.1 1.9 3.46 0.79 104 2.2 2.89 1.20 180 2.3 1.87 1.41 288 2.6 3.03 2.24 285 2.6 4.71 0.92 107 1.8 12.3 2 .o 72.6 2.2 11.4 2.6 102 2.2 8.8 4.4 234 2.1
(c) Flowtube Freshly Coated with Halocarbon Wax 22.4 20.4 2 .o 233 0.42 23.2 20.2 3 .O 187 0.79 22.4 21.4 1 .o 137 0.34 17.5 46.8 0.7 148 0.28 17.2 14.9 2.3 221 0.70 16.8 15.1 1.7 263 0.43 19.0 17.6 1.4 124 0.64 18.4 17.8 0.6 183 0.18 17.4 15.4 2 .o 258 0.47 17.0 14.6 2.4 318 0.52 17.9 15.3 2.6 226 0.75 15.1 14.1 1 .o 138 0.51 14.1 13.1 1 .o 215 0.36 13.7 12.9 0.8 250 0.25 14.4 14.1 0.3 191 0.11 11.9 11.6 0.3 172 0.15 11.3 10.3 1 .o 262 0.37 12.3 11.8 0.5 103 0.41 (d) Measurements Made on Aged Halocarbon Wax
4.60 4.25 4.66 5.42 6.11 4.19 a
3.14 2.76 3.54 3.71 5.16 2.89
1.46 1.49 1.12 1.71 0.95 1.30
208 349 164 304 113 193
2.2 1.6 1.9 1.5 1.6 2.3
Concentrations in units of 1013 molecule/cm3.
cm3 molecule-', the value calculated for [ ( c l ) ~(Cl)/(HCl)(Cl)] in these tubes. It was changed to 4.3 x for experiments performed on fresh halocarbon wax. (c) k5 was allowed to vary between and cm3 molecule-l sec-l, which is the smallest value that can be chosen if we insist upon establishing equilibrium between C1 and HCl within 3 msec. All of the results shown in Table IV were obtained with k g equal to 1.5 X 10-l' cm3 molecule-' sec-l, and k-5 set equal to 2500 sec-l. A good fit with experiment was obtained, and the following features were accurately reproduced: (a) the acceleration in hydrogen atom decay brought
(a) 100 100 (b) 100 100 (c) 106 106 (d) 100 100
(e) 64 to 135 (f) 64 to 135 (g) 107 to 122
(H),
(Cl),
10 0 10 10 5 0 5 5 18.9 0 12.0 6.9 10 0 5 0 1 to 0 19 1 to 6.9-7.9 19 3 to 0 22
k,/-k,
2.1 x 2.1 x 2.1 x 2.1 x 2.1 x 2.1 x 4.3 x 4.3 x
0.48 0.32
8.1 18.0 0.59 0.44 7.1 10-'6 14 .O 0.37 0.21 9.0 10-16 21.0 lo-'? 0.83 0.72 5.9 lomi70.88 0.81 5.3 10-16
2.1 x 2.1
X
4.3
X
At 5 At 10 msec msec k,,,
0.33 0.29
7.3 12.3
lo-'' 0.70 0.77
5.8
a The rate constants used are given in the text. k , , , given in units of 10-14 cm3 molecule-l sec-I. Lines e, f, and g are experimental averages extracted from Tables I and I1 for comparison purposes. No variation with stoichiometry was observed experimentally. Concentrations are in units of 1013 molecule/cm3.
about the addition of extra atomic chlorine; (b) the increase in the rate of hydrogen atom removal that occurred on changing the flowtube coating from fresh halocarbon wax, with K = 4.3 X lo-", to fluorinated halocarbon wax, with K = 2.1 X and (c) the decrease in the ratio of C1 formed to H removed that was observed on changing the flowtube coating. It should be noted that this fit with experiment was obtained while preserving the equality between kJk-1 and the equilibrium constant. Unfortunately the mechanism outlined above is an artificial structure chosen mainly for its simplicity, and for its ability to model certain rate parameters. In the following sections questions concerning the real chemical validity and significance of the model are discussed. Evidence for the Stability of HC12. No unambiguous evidence for the existence of HClz exists in the literature. However, several reports suggest that the structure Cl-HC1 may be stable. Truhlar, Olsen, and P a d 5 examined ClHCl using the bond-order-bond-energy (BEBO) method of Johnston and Parr, and found it to be bound and stable. They estimated the depth of its potential energy minima as 1.56 kcal/mol, and its equilibrium internuclear H-Cl distance as 1.45 A. Although theorists tend to be skeptical of empirical schemes such as BEBO, the method does predict a minimum energy path for the F H2 reaction that is within 0.03 A of that obtained from the most elaborate ab initio calculation,16 and thus it should be considered seriously when applied to systems that are chemically similar. More evidence for the existence of ClHCl was provided by Noble, and Pimental,17 who reported its matrix-isolated infrared spectrum. These authors added HCl to the products of an electric discharge in a dilute ClZ-Ar mixture, and condensed the resulting gases a t 20 K. The infrared spectrum of this matrix contained prominent absorptions at 956 and 656 cm-l, and these were assigned to the linear, symmetric structure, Cl-H-Cl. Isotopic substitution tended to confirm this assignment, and the stability of ClHCl was estimated a t 5-10 kcal/mol. Unfortunately, this seem-
+
The Journal of Physical Chemistry, Vol. 79, No.22, 1975
2334
John E. Spencer and G. P. Glass
ingly unambiguous evidence has been challenged by Milligan and Jacox18 who claim that the molecule isolated was, in fact, the negative ion, ClHC1-. This claim has some merit, but, at this time, no definitive experiment allowing us to choose between the two assignments has been made. Some additional evidence for the stability of HClz was obtained in this study by measuring the line broadening cross section for collision of C1(2P3/a)with HCI. Cook and Miller19 have shown that these cross sections for P state atoms are determined largely by long-range attractive forces. Therefore, the line broadening cross section for the collision between C1 and HCI provides a measure of the magnitude of the long-range forces between these species. In this study, the cross section was estimated from EPR line width measurements. Under conditions of low magnetic field modulation, and low microwave power, these are related by the equation (AH)-' = ( 1 2 / ~ k T ) ~ / ~ 6 - ' ~~ui/y,'/~ P 1
(9)
0.72
1-
0.48l
I
0.02
I
I
0.06
I
I
0.10
I
1
0.14
MOLE FRACTION HCI Figure 4. The effectof HCI on the EPR line width of atomic chlorine. The line observed is for the transition C1(2P3/2) [Mi = 3/2,MJ,= -'A -t- %].Total pressure kept constant at 0.72 Torr. The major constituent, other than HCI, was argon.
-
where 6 is the gyromagnetic ratio, P the total pressure, x, the mole fraction of the ith component, CT& the line broadAlthough HC1 and C1 must partly react at the flowtube ening cross section for collision between the ith component walls, the resulting product must be liberated into the gas and the absorbing species, and y, the reduced mass for this phase, as the followingdiscussion shows. Consider the chlocollision. rine atom flux across a section of the flowtube at two posiIn our experiments, the EPR spectrum of 35C1(2P3/2,mI = -3/2, mJ = -l/2 M) was recorded as HCl was added to tions, one upstream, and the other downstream from the various dilute C1-Ar mixtures. The EPR line width, AH, HC1 inlet. Between these positions, the flux decreases. Therefore, the chlorine atom reaction product must be ( 1 ) was plotted against the mole fraction of added HC1, and continuously deposited on the walls between these posithe line broadening cross section was estimated from the tions, or ( 2 ) liberated into the gas phase. If the first possigradient. Seven experiments were performed, each similar bility occurred, the walls would soon become coated with to the one shown in Figure 4. The cross section UCI-HCI was the product molecule, its thermodynamic activity would estimated as 120 f 14 X cm2. Westenberg20 has meaapproach unity, and (Cl)(HCI) would approach a constant sured the cross section for collision of C1 with Ar, which is value, analogous to that of the solubility product for liquid/ cm2. This value has isoelectronic with HCl, as 32 X been recently corroborated by Zegarski, Cook, and Miller,21 solid systems. This was not observed. Therefore, we favor the following mechanism for the wall formation of HC12: and agrees well with that estimated in this work (Figure 4 at HCl = 0). Thus, the magnitude of the cross section for C1+ HCl HClz(wal1) (10) collision with HC1 suggests that strong long-range attractive forces exist between C1 and HCI. Possible Mechanisms for the Formation of HCIZ. The A major objection to this mechanism is that it requires data presented in Tables I and I11 clearly show that the the measured ratio [(Cl)0 - (CI)]/(HCl)(Cl) to be indepenrate of overall reaction, and the diminution in atomic chlodent of wall coating. This requirement arises because, at rine concentration upon addition of HC1, both depend on equilibrium, the ratio is equal to the equilibrium constant the condition of the flowtube coating. This strongly for the process suggests that the formation of HCl2 largely occurs at the flowtube walls. However, a RRKM calculation of the life(12) H(g) + HCUg) == HCMg) time of HC12, formed in the gas phase by collision of C1 and and this is determined solely by the standard free energy HCI, indicates that some HClz may be formed homogedifference between these gas phase species. At first sight, neously. In this calculation, the structure of HC12 was asthe ratios listed in Table I11 do not appear to be indepensumed to be that proposed by Truhlar et aI.,l5 its dissociadent of wall coating. However, the low group of values were tion energy was chosen as 10 kcal/mol, and its excess enermeasured using a freshly prepared halocarbon wax coating, gy of formation as 600 cal/mol (RT). The lifetime, estimatand it is possible that these values represent a steady state ed as approximately IO-'' sec, did not much vary as the ascondition not representative of true equilibrium. The sumed structure of HClz was altered, and was largely deterchange in the ratio, observed after these surfaces had aged, mined by the total number of atoms in the adduct. At the supports this contention. pressures used in this study (-1 Torr), collisional deactivaThe equilibrium constant for reaction 12 is related to the tion of HCl2 within its estimated lifetime occurs with a potential energy of dissociation of HC12, A E , by the equaprobability of approximately Therefore, the predicttion ed rate of the homogeneous reaction corresponds to a value for k 5 of 3 X cm3 molecule-' sec-'. This is a factor of K s, [Z(HCIZ)/Z(HC~)Z(CI)] exp(AE/RT) 3 below what is needed to explain the rapid attainment of where Z(HC12) is the partition function for the HC12 moleequilibrium between C1 and HC1 shown in Figure 3. Howcule. If the equilibrium constant is set equal to 2.2 X 10-16 ever, since very weak ArHCl complexes have been obcm3 molecule-', the average value of [ ( c l ) ~ (CY)]/ served,22homogeneous formation of HC12 might also have a (HCl)(Cl) measured on surfaces coated with aged or fluosmall contribution from the "bound complex" mechanism.
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The Journal of Physical Chemistry, Vol. 79, No. 22, 1975
Electronic Spectra of Thioacetic Acid
2335
rinated halocarbon wax and on surfaces washed with 5% HF, and if the structure of HClz is assumed to be that proposed by Truhlar et a1.,'6 then the potential energy of dissociation of HCl2 can be estimated as 9.8 kcal/mol. This value is not affected greatly by changes in the structure of HC12, but is, of course, totally dependent upon the correctness of the assumed HC12 formation mechanism.
flowtubes coated with boric acid. In such tubes, wall recombination of atomic chlorine occurs, and thus, their results may have been influenced by the reaction sequence:
Conclusions The major conclusion reached in this study is that the ratio of the thermal rate constants kJk-1 for the system H HCl H2 C1 is equal to the equilibrium constant of this system. In our experiments, we find that atomic chlorine, produced in the reaction H HCl HP C1, reacts with HCl, largely at the wall, to form a product (possibly HClZ), that reacts further to consume another hydrogen atom. The apparent rate constant estimated from the rate of disappearance of atomic hydrogen in the presence of excess HCl is, therefore, much greater than the elementary rate constant 121. A t 298 K, kl can be estimated from the equilibrium constant and the back reaction as 4.3 X cm3 molecule-' sec-l. This value is also obtained if our results are extrapolated to a ratio for C1 formed to H removed of unity, a condition implying the absence of secondary reactions. Direct wall reactions between H and C1 do not occur to any significant extent. However, the secondary reactions outlined above do produce effects similar to those produced by the mechanism postulated by Galante and Gisla-
Acknowledgments. Acknowledgment is made to the Donors of the Petroleum Research Fund, administered by the American Chemical Society, for the partial support of this work. One of us (J.E.S.) thanks the National Science Foundation for a Fellowship.
+
+
+
-
+
The earlier measurements of Westenberg and deHaasl were undoubtably affected by the secondary reactions of atomic chlorine described above. In particular, the increase in the H atom decay rate, observed as the absolute H atom concentration was raised, is well explained by a mechanism involving HC12. This is illustrated clearly in Table IV. However, their experiments were performed using
c1 H
+ C12
-
wall
-+
Cl2
HCl
+ C1
References and Notes ( I ) A. A. Westenberg and N. deHaas, J. Chem. Phys., 48, 4405 (1968). (2)N. S.Snider, J. Chem. Phys., 53, 41 16 (1970). (3)M. A. A. Clyne in "Physical Chemistry of Fast Reactions", Vol. 1, B. P. Levitt, Ed., Plenum Press, New York, N.Y., 1973,p 245. (4) J. J. Galante and E. A. Gislason, Chem. Phys. Lett., 16, 231 (1973). (5)D. D. Davis, W. Braun, and A. M. Bass, Int. J. Chem. Kinet., 2, 101 (1970). (6)S . W. Benson, F. R. Cruickshank, and R. Shaw, Int. J. Chem. Kinet., 1, 29 (1969). (7)J. E. Breen and G. P. Glass, J. Chem. Phys., 52, 1082 (1970). (8) G. A. Takacs and G. P. Glass, J. Phys. Chem., 77, 1060 (1973). (9)A. A. Westenberg, Prog. React. Kinet., 7 , 24 (1973). (IO) W. A. Breckenridge and T. A. Miller, J. Chem. Phys., 56,475(1972). (1 1) Halocarbon Products Corp, 82 Burlews Court, Hackensack, N.J. (12)D. H. Stedman, D. Steffenson, and H. Niki, Chem. Phys. Lett., 7 , 173 (1970). (13)D. Arnoidi and J. Wolfrum, Chem. Phys. Lett., 24, 234 (1974). (14)M. A. A. Clyne and D. H. Stedman, Trans. Faraday SOC., 62, 2164 (1966). (15)D. G. Truhlar, P. C. Olsen, and C. A. Parr, J. Chem. Phys., 57, 4479 (1972). (16)D. G. Truhlar, J. Am. Chem. SOC.,94, 7584 (1972). (171 P. N. Noble and G. C. Pimental, J. Chem. Phys., 49,3165 (1968);55, 540 (1971). (18) D. E. Miiligan and M.E. Jacox, J. Chem. Phys.. 53, 2034 (1970). (19)T. A. Cook and T. A. Miller, J. Chem. Phys., 99, 1352 (1973). (20)A. A. Westenberg and N. deHaas, J. Chem. Phys., 51, 5215 (1969). (21)6. R. Zegarski, T. J. Cook, and T. A. Miller, J. Chem. Phys., 62, 2952 (1975). (22)S. E. Novick, P. Davies, S.J. Harris, and W. Klemperer, J. Chem. Phys., 59, 2273 (1973).
A Study of the Electronic Spectra of Thioacetic Acid and Its Ethyl Ester Shinichi Nagata; Tokio Yamabe, and Kenichi Fukui Department of Hydrocarbon Chemistry, Faculty of Engineering, Kyoto University, Sakyo-ku, Kyoto, Japan (Received December 13 1974; Revised Manuscript Received July 21, 1975)
Electronic spectra of thioacetic acid (TAA) and its ethyl ester (ETA) were measured in ethanol and cyclohexane. Their absorption bands were characterized by their small excitation energies as compared with those of the analogous compounds, CH3COOH or CH3CONHz. These bands were reasonably assigned by means of an ASMO-SCF-CI calculation. Furthermore, in order to clarify the relationship between ground and excited states, we measured the photoelectron spectra of TAA and ETA, which provided valuable information for the interpretation of the characteristic bands of T A A and ETA.
It is well known that organic sulfur compounds often have remarkable difference in valence,' bond length, and bond angle in comparison with similar compounds contain-
ing oxygen atoms in place of the sulfur, and also show different reactivities or even unexpected reactions. Recently there have been a number of theoretical investigationq2 The Journalof Physical Chemistry, Vol, 79, No. 22, 1975
