THE
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PHYSICAL CHEMISTRY
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@ Copyright, 1971, by the A m e r i m n Chemical Society
VOLUME 75, NUMBER 12 JUNE 10, 1971
Mass Spectral Study of the Decomposition of Chlorine Fluoride behind Shock Waves1 by J. A. McIntyre* and R. W. Diesen Chemical Physics Research Laboratory, The Dow Chemical Company, Midland, Michigan 48640 (Received November 4, 1970) Publication costs assisted by T h e Dour Chemical Company
The thermal decomposition of chlorine fluoride has been studied over the temperature range 2000-2950'K using the technique of mass spectral sampling behind reflected shock waves. The reaction has been found to be complex, and evidence is offered to support the suggested mechanism which consists of the reactions
+F+M C1 + C1F J 'Clz + F
CIF
+M
A C1 kz
k-a
From experimental data, k-z is estimated to be greater than 1 X 1012 cc/mol sec.
+
Introduction Fz RI +2F It is surprising to note that, of the very few k n o ~ n ~ - I~n eq 1-5, M is a collision partner.
investigations of the thermal decomposition of the chlorine fluorides (ClF, C1F3,ClFs), the simplest system, CIF, for which an unambiguous interpretation of experimental results is most likely, was most recently undertal~en.~Even in this system the kinetics may be complex, since many reactions are possible at high temperatures in a dissociating system containing two or more atomic species. The reactions listed below are those that seem most likely to be important under the conditions of this investigation.6
+ 34 -+ C1 + F + 34 C1 + C1F J_ GIz + F
ClF
Clz
F
+ M * 2C1 + M
+ C1F
C1
+ Fz
(1)
(3)
(4)
+ 14
(5)
Some of the above reactions have been suggested' as being important steps in the mechanism of the reaction between C12 and F2 to yield C1F. That study, performed in a static reactor a t temperatures between (1) This work was sponsored by the U. 8. Office of Naval Research under Contract No. Nonr 3814(00). (2) J. A. Blauer, H. G. McMath, and F. C. Jaye, J . P h y s . Chem., 73, 2683 (1969). (3) A. E. Axworthy and J. M.Sullivan, ibid.,74, 949 (1970). (4) J. A. Blauer, H. G. McMath, F. C. Jaye, and V. S. Engleman, ibid., 74, 1183 (1970). (5) J. A. Blauer, V. S. Engleman, and W. C. Solomon, private communication. (6) For Clz and Fz-free ClF, reactions 3 and 6 can only be important if reactions 2 and 4 are. (7) E. A. Fletcher and B. E. Dahneke, J . A m e r . Chem. SOC.,91, 1603 (1969).
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J. A. MCINTYRE AND R. W. DIESEN
130 and 170") disagreed with theoretical predictions which had been mades*O(and which also disagreed) concerning that reaction. With this backgroundj6 a study of the thermal decomposition of chlorine fluoride by means of the technique of mass spectral sampling behind reflected shock waves was undertaken. For dilute solutions of C1F in neon in the temperature range 200O-295O0K, we have found the reaction to be complex and report here a mechanism which adequately accounts for the known experimental results. Experimental Section The apparatus and the methods used to obtain kinetic information have been described previously.10-12 Samples were prepared from Ozark-RIahoning ClF, purifiedI3 just before use by trap-to-trap distillation at various temperatures between - 80 and - 145") Matheson research grade Clz (also distilled), General Chemical Fz (99%)) Dow argon (99.99%)) and Matheson research grade neon. Airco helium and Dow argon were used as driver gases without purification. For all samples neon was used as the primary diluent, with argon added at 0.25% as an internal pressure standard.'O Mixtures for decomposition experiments contained nominally14 0.5% reactant. Mixtures for the Clz Fz reaction studies contained 0.5% of each, and were ClF-free.
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Results and Discussion By analogy with the halogens, one may be inclined to treat the experimental results for the disappearance of C1F in terms of the pseudo-first-order reaction (1) above, where R4 is the inert gas diluent. Using this analysis we have obtained apparent rate constants over the temperature range 2000-2950°K. An Arrhenius plot of these apparent rate constantsI6 is shown in Figure 1. From a least-squares treatment of the data the apparent rate constant, k,, is described by the equation
Our experimental activation energy is in reasonable agreement with that normally observed for Clz dissociation in shock t ~ b e s . ~ ~ ' Since ~ - * ~ the dissociation energies of Clz and ClF only differ by approximately 2% (1.2 kcal/mol), this agreement would not be unexpected, particularly since the activation energies for halogen dissociation generally have been found (see, for instance, ref 20 for discussion) to be lower than the bond dissociation energy. However, additional information, discussed below, indicates that the reaction is more complex than this. For the C1F system, the mass spectral technique gives one the advantage of being able to monitor independently and simultaneously the disappearance of C1E' and The Journal of Physical Chemistry, Vol. 76, N o . 13, 1971
-
lo"
\oo
I
I
I
0
I
~~
0
lo9
O
\
X
lo8
I
Figure 1. Arrhenius plot of apparent rate constants: 0, first-order rate constants for the disappearance of ClF in nominally pure ClF samples, based on 48 experimental values; X, first-order rate constants for t h e disappearance of Clz in ClF mixtures. Clz
+
Clz. Thus, when the data for the disappearance of Clzin mixtures containing small amount'sof C1F are interpreted (8) S. W.Benson and G. W. Haugen, J . Amer. Chem. Soc., 87, 4036 (1965). (9) R. M. Noyes, ibid., 88, 4311, 4318 (1966). (10) R. W. Diesen and W. J. Felmlee, J . Chem. Phys., 39, 2115 (1963) (11) R. W,Diesen, ibid., 39, 2121 (1963). (12) R. W. Diesen, ibid., 44, 3662 (1966). (13) Considerable care and effort were spent in attempts to obtain 100% C1F; mass spectral analysis always showed 1-2% residual Ch. We remain skeptical of claims of 100% purity due to the extreme reactivity of C1F. We have found that the rate of disappearance of C1F increases rapidly in the presence of small amounts of oxygen. The mass spectral sampling technique allows us to exclude this possibility for the results we report here. (14) The purity of C1F has been discussed.*3 Our own attempts a t preparing CIF gave a product high in C12, which was used for the CL in C1F experiments. (15) For reactions which will become obvious we do not tabulate these rate constants. The reflected shock pressures were in the range 300-550 Torr. (16) H. Hiraoka and R. Hardwick, J . Chem. Phys., 36, 1715 (1962). (17) T . A. Jacobs and R. R. Giedt, ibid., 37,747 (1963). (18) M.van Thiel, D. J. Seery, and D. J. Britton, J. Phys. Chem., 69, 834 (1965). (19) R. A. Carabetta and H. B. Palmer, J . Chern. Phys., 46, 1333 (1967); 47, 2202 (1967). (20) See, for example, H. 0. Pritchard, Transfer Stor. Energy Mol., 2, 368 (1969); G. Burns and R. J. Brown, J . Chem. Phys., 53, 3318 (1970). I
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DECOMPOSITION OF CHLORINE FLUORIDE in terms of reaction 3, the apparent rate constants are found to be very similar t o those resulting from an equivalent treatment of the data for nominally pure C1F (see Figure 1). Further, the apparent rate for the disappearance of Clz in nominally pure CIF (where C12 is about 3-501, of the CIF concentration) is found to be very small, and in some cases C12increases slightly. Moreover, in the same shock, the apparent rate constants for Clz disappearance became comparable. Thus, the results imply coupling between the C1F and Clzand suggest the importance2’of reaction 2. I n this case, reaction 3 also must be considered. This leads then to the suggested mechanism, which involves the reactions (rewritten here for convenience)
+M C1 + C1F Clz + NI
C1F
+F+M Cl2 + F 2c1 + M
(1)
C1 k2 k-2
(2)
(3)
From these equations one obtains -d[ClF] dt
~-d[C1zl
ki[M] [ClF]
-
+
k3[M][Cl2] -
dt IC2
[GI][ClF]
+
k-2
[F
or, from addition of (7) and (8) -d[ClF] dt
+---d[Clz] dt
kl [MI [ClF]
+ kg [M
which can be rearranged to
-din [CIF] dt
Note that eq 10 involves no approximations or assumptions other than that reactions 1, 2, and 3 describe the mechanism. To have any validity the proposed mechanism must account for the observations that (a) the disappearance of C1F can be described by pseudo-firstorder kinetics; (b) at 1o;ver temperatures in high [ClF]/ [Clz] samples, the [Clz]remains constant or slightly increases, while at high temperatures it decreases a t a rate which is (apparently) described by a pseudo-firstorder plot with rate constants similar to those for ClF. Point (a) is accounted for as follows. At high temperatures the experimental quantitites - d In [C12]/dt (ie., the slope of a first-order plot) and the ratio [C12]/ [ClF] remain essentially constant. This means that
-d In [ClF]/dt will be essentially constant (from eq lo). At lower temperatures, [Cl:,] is observed t o be essentially constant (which means -d In [Clg]/dt is very small), and again -d In [ClF]/dt appears to be constant. We do not claim that -d In [ClF]/dt is truly constant-only that it would require rate data of very high precision in order to see the small degree of curvature.22 The intent here is to demonstrate how CIF can apparently disappear by a first-order reaction even if the equation governing its disappearance is more complex. The experimental observations of the fate of Clz in this system make it necessary to invoke reactions other than (1). For reaction 2 to make significant contribution, kz must have a large value. For reaction 2 to be important in the early stages of the reaction, kz must be orders of magnitude larger than kl since at these times species other than C1F will be at low concentration. Experiments on the reaction of C1, and Fz in the shock tube over a similar temperature range (1450-27OO0K), show that CIF is produced rapidly (which is not unexpected), and a distinct induction period is observed. The results are consistent with the interpretation of the low-temperature static work.’ Experimentally the induction period for the formation of C1F decreases with increasing temperature. It can be shown thatz3 At 0: (1/k-2[C1z]) where At = induction period. From this expression we estimate that k-z is greater than 1 X 10l2 cc/mol sec (e.y., at mol/cc, At -50 X 2040°K, [C12] = 1.57 X sec). The value of can reasonably be taken to be a lower limit since the proportionality constant will be larger than unity. This value of k-2 is certainly large enough to make reaction 2 important. Since the equilibrium constant2*for reaction 2 is about 0.6 in this temperature range, k2 is estimated to be 21 X 10l2 cc/mol sec. Thus, k2 is certainly large enough to make reaction 2 important in the C1F decomposition. This large value of the rate constant, in conjunction with an assumed low activation energy for reaction 2, further explains the results. 25 At high temperatures, dissociation (ie., disappearance of total molecules) occurs by reactions 1 and 3, but these rates are coupled through the fast reaction 2. At lower temperatures reactions (21) The absence of Fz in the reacting mixture precludes reaction 4 (and thus 5) being important. (22) Over one order of magnitude, allowing some scatter in the data, one can fit such data equally well by In z or In (a z) = kt, where a < x. (23) See, for instance, S. W. Benson, J . Chem. Phys., 20, 1805 (1952). (24) “JANAF Thermochemical Tables,” The Daw Chemical Co., Midland, Mich., June 1970. (25) For supporting evidence see K . H. Homann, J. Warnatz, H. Gg. Wagner, and C.Zetsch, Ber. Bunsenges Phys. Chem., to be published. See also A. Hoffman, K. H . Homann, and D . I. MacLean, ONR Report FRK-111, Boston College, Chestnut Hill, Mass., Jan 1971.
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The Journal of Physical Chemistry, Val. 76, ATa. 12, 1971
1768
K. KIMURA, T. YAMAZAKI, AND S. KATSUMATA
1 and 3 become less important relative t o reaction 2 due to their higher activation energies. Thus, in our samples (large [ClF]/ [Clt]), the CIF still disappears (via reaction 2) but now the [Clz] remains constant or even increases slightly depending on the actual temperature.
Conclusions The thermal decomposition of C1F is complex in the temperature range 2000-2950°K, and a mechanism which adequately explains the experimental results is suggested. It is reemphasizedz6that caution is in order
when inferring mechanisms from semilog kinetic plots. The results reported here also indicate that k, and k3 are similar in magnitude although a more detailed study would be required t o evaluate IC1 explicitly.
Acknowledgment. The authors wish to thank R. W. Anderson and H. R. Frick for the synthesis of ClF, and W. E. Eichorn for his help in obtaining some of the experimental data. (26) A different way in which aroblems can arise has been demonstrated recently: M. W. Slack and E. s. Fishburne, J . Chem. P h i s . , 52, 5830 (1970).
Dimerization of the Perylene and Tetracene Radical Cations and Electronic Absorption Spectra of Their Dimers by Katsumi Kimura,* Tomoko Yamazaki, and Shunji Katsumata Physical Chemistry Laboratory, Institute of Applied Electricity, Hokkaido University, Sapporo, J a p a n (Received October 6 , 1970) Publication costs assisted by the Institute of Applied Electricity, Hokkaido University
The reversible dimerizations of the perylene and tetracene radical cations in concentrated sulfuric acid have been studied spectroscopically. Equilibrium constants of the monomer-dimer equilibrium at various temperatures were determined by analyzing electronic absorption spectra. From the temperature dependences of the equilibrium constants, heats of dimerization were evaluated to be 8.8 and 5.6 kcal/mol of dimer for the perylene and tetracene cations, respectively. From the present study it has also been found that when the cation dimers are formed new absorption bands appear in the near-infrared region, and the ultraviolet and visible absorption bands shift toward shorter wavelength. The appearance of the near-infrared bands in the dimer spectra are discussed in terms of both a radical-radical charge resonance and an electronic transition between the bonding and antibonding MO’s constructed for the dimer, on the basis of a symmetrical sandwich dimer structure. Interactions between the two cation radicals forming the dimer in solution are discussed.
Introduction Recently, a considerable amount of study has been reported on the dimerization of organic free radical (neutra11i2 or ionic1v3-13) in solutions. Most of the radical ions whose dimerizations have been studied so far are for substituted benzenes, while very little is known on dimerization of the radical cations and anions of catacondensed hydrocarbons. For the perylene and tetracene cations (perylene+ and tetracene+) it has now been found from electronic absorption studies that the dimers of these radical cations exist in equilibrium with the monomers in concentrated sulfuric acid solutions and can be detected even a t concentrations as M a t room temperature. low as about I n previous papers, Kimura, et al., have studied dimerization phenomenon of the p-phenylenediamine catThe Journal of Physical Chemistry, Vol. 76, No. 18,1971
ion in solution at low temperature by analyzing electronic absorption spectra quantitatively4 and have (1) K. H. Hausser and 3 . N. Murrell, J . Chem. Phys., 27, 500 (1957). (2) (a) M.Itoh and S. Nagakura, J . Amer. Chem. SOC.,89, 3959 (1967); (b) D. A. Wiersma and J. Kommandeur, Mol. Phys., 13, 241 (1967); (0) K . Maeda and T . Hayashi, Bull. Chem. SOC.Jap., 42, 3509 (1969). (3! K. Uemura, 5 . Nakayama, Y. Seo, K. Suzuki, and Y. Ooshika, abPd., 39, 1348 (1966). (4) K. Kimura, H. Yamada, and H. Tsubomura, J . Chem. Phys., 48, 440 (1968). (5) H. Yamada and K. Kimura, ibid., 51, 5733 (1969). (6) T. Sakata and S. Nagakura, Bull. Chem. SOC.J u p . , 42, 1497 (1969). (7) J. Tanaka and M . Mizuno, ibid., 42, 1841 (1969). (8) G. Spach, H . Monterio, M . Levy, and M . Szarc, T r a n s . Faraday SOC.,58, 1809 (1962). (9) E. M .Kosower and J. L. Cotter, J . Amer. Chem. Soc., 86, 5524 (1964).
