J. Phys. Chem. 1986, 90, 1988-1990
1988
if the detailed nature of such stratospheric surfaces were firmly established. Both reactions -9 and 10 form Cl20, a molecule not normally included in present atmospheric models. While reaction 9 should furnish a rapid removal route for Cl2O in the atmosphere, other exothermic reactions competitive with it may also exist for C120. The reaction of CI20with HCI by (1 1) is a known rapid reaction and would tend to convert the C1 trapped in relatively inert HCI into an easily photolyzable form. Exothermic reactions can also whose stratospheric concentration is be found for ClzO with 03, comparable to that of H20. One very important attribute of the heterogeneous oxide exchange reactions is that they do not require sunlight and can occur equally well at night. This could result in substantial shifts in chlorine distributions from those expected during the daylight periods dominated by free radical reactions driven by solar radiation. The total organochlorine concentration of the troposphere has increased by a factor of three in the past 15 years to a 1986 level of about 3500 pptv (l0-l2) in the northern hemisphere and is continuing to increase at a rate of at least 1000 pptv per decade.'
The recent rapid change in the yearly decrement in total ozone in the Antarctic ~ p r i n g ~appears -~ to correspond more to a quadratic than to a first-power dependence on the total tropospheric chlorine content. Several of the reactions, e.g. (7), (lo), and (1 l ) , involve the interaction of two chlorinated species and could, if fast enough, contribute to a higher than first-order dependence on total organochlorine concentration. Reasonable simulations of the Antarctic ozone observations have recently been attained by inclusion in an atmospheric model of heterogeneous reactions of C 1 0 N 0 2 with either HCl or H20.26
Acknowledgment. This research was supported initially by the Office of Basic Energy Sciences of the Department of Energy through Contract No. DE-AT-03-76ER-70126, and then by NASA Contract NAGW-668. The authors thank Drs. Atkinson and Winer for discussions and a preprint of their work on chlorine nitrate. (26) Solomon, S.;Garcia, R.; Rowland, F. S.; Wuebbles, D. Nature (London), in press.
Reaction Rates for Thermal Chlorine Atoms wlth H,S from 232 to 359 K by a Radiochemical Technlque Eric C. C. Lu, R. Subramonia Iyer, and F. S. Rowland* Department of Chemistry, University of California, Irvine, California 9271 7 (Received: December 3, 1985; In Final Form: January 28, 1986)
The relative rate constants for the gas-phase reactions of thermal chlorine atoms have been measured with H2S, C2Hs, and CH2=CHBr over the temperature range from 232 to 359 K by using radioactive 38Clatoms. These data at a pressure of 4000 Torr of CClF, are consistent with a temperature-dependent rate constant of (10.5 0.4)X lo-" cm3 molecule-' s-l, based on the consensus literature values for the reaction of C1 with C&.
*
We have measured the reaction rate constants for chlorine atoms with H2S by (1) over the temperature range from 232 to 3sCl
+ H2S
-
+ SH
H38C1
(1)
359 K utilizing a competitive radiochemical technique at pressures of approximately 4000 Torr.I4 The reaction of chlorine atoms with H2S is potentially important in the atmosphere of the earth and of other planets.56 Several previous experiments have provided measurements of the reaction rate constant at 298 K ranging from 4 X lo-" to >7.3 X lo-" cm3 molecule-' s-IS5-l1 The one study over a temperature range has indicated a rate constant of (6.29 f 0.46) X lo-" cm3 molecule-' s-l from 21 1 to 353 K at pressures (1) Lee, F. S. C.; Rowland, F. S.J . Phys. Chem. 1977, 81, 1235. (2) Lee, F. S. C.; Rowland, F. S. J. Phys. Chem. 1977,81, 1229. (3) Lee, F. S. C.; Rowland, F. S. J . Phys. Chem. 1977, 81, 86. (4) Lee, F. S. C.; Rowland, F. S. J . Phys. Chem. 1980, 84, 1876. (5) DeMore, W. B.,Ed. "Chemical Kinetics and Photochemical Data for Use in Stratospheric Modeling. Evaluation Number 7"; Jet Propulsion Laboratory: Pasadena, CA, 1985; JPL-85-37. (6) Baulch, D. L.; Cox, R. A.; Crutzen, P. J.; Hampson, R. F.; Kerr, J.; Troe, J.; Watson, R. T. J. Phys. Chem. Ref.Data 1982, 11, 327. (7) Braithwaite, M.; Leone, S. R. J . Chem. Phys. 1978, 69, 839. (8) Nesbitt, D. J.; Leone, S. R. J. Chem. Phys. 1980, 72, 1722. (9) Clyne, M. A. A.; Ono, Y.Chem. Phys. Lett. 1983, 94, 597. (10) Clyne, M. A. A.; MacRobert, A. J.; Murrells, T. P.; Stief, L. J. J. Chem. SOC.,Faraday Trans. 2 1984, 80, 877. (11) Nava, D. F.; Brobst, W. D.; Stief, L. J. J . Phys. Chem. 1985, 89, 4703.
0022-3654/86/2090-1988$01.50/0
of 100 Torr of Ar or less.'' Our data at 4000 Torr of CCIF, are consistent with a temperature-independent rate constant of (10.5 f 0.4) X lo-" cm3 molecule-' s-l from 232 to 359 K. Radioactive 38Clatoms have been formed by thermal neutron irradiation of CClF3 and moderated to thermal energies by multiple collisions with CCIF3 prior to competitive reactions with one of two minor substrate molecules.'-" The basic competition has been between reaction 1 with H2S and reaction 2 with 38CI + CH2=CHBr
-
CH2=CH38Cl
+ Br
(2)
CH2=CHBr, the latter leading to the formation of the easily measured molecule CH2=CH3sCl.'2 Both reactions have been placed on an absolute reaction rate scale through separate study of the competition between (2) and the well-known abstraction reaction of H from C2H6 in (3).5*6 The monitor reaction with
+
38Cl CzH6
-
H38Cl
+ C2H5
(3)
CH2==CHBr in (2) has been chosen because it is approximately as fast as its two competitor reactions (1) and (3) and because it produces a substantial yield of CH2=CH38C1 which can be readily separated and assayed by radio gas chromat~graphy.'~-'~ (12) Iyer, R. S.; Rowland, F. S. Chem. Phys. Lett. 1983, 103, 213. (13) Iyer, R. S.; Rogers, P. J.; Rowland, F. S.J . Phys. Chem. 1983, 87, 3799. (14) Lee, J. K.; Lee, E. K. C.; Musgrave, B.; Tang, Y.-N.; Root, J. W.; Rowland, F. S. Anal. Chem. 1962, 34, 741.
0 1986 American Chemical Society
Letters YCH2.CH38CI 0.95
15
/
The Journal of Physical Chemistry, Vol. 90, No. 10, 1986 1989
0.95
0.95 YCH2.CH38CI
YCH2=CH38CI
232 K
15
IO
359 K
10
5
C
1
2
3
4
5
(RH) /(CH2=CHBr)
“
0
1
2
3
4
5
(RH) / (CH2=CHBr1
0
0
1
2
3
4
5
(RH) /(CH2=CHBr)
Figure 1. Reciprocal yields of CH2=CH38C1vs. [H2S]/[CH2=CHBr] or [C2H6]/[CH2=CHBr]: left, 232 K; center, 295 K; right, 359
* 4 K.
TABLE I: Relative and Absolute Rate Constant Measurements for Chlorine Atom Reactions with H S and C2Hn in Competition with CH,=CHBr
temp, K 232 295 359 f 4
k,lk2 2.95 f 0.13 2.24 f 0.07 0.83 f 0.04
kdk2 1.48 f 0.06 1.20 f 0.06 0.47 f 0.02
kdk3 1.99 f 0.12 1.87 f 0.11 1.77 f 0.11
10”k3, cm3 molecule-’ s-I [5.221 [5.68j [5.99]
10”kl, cm3 molecule-’ s-I 10.4 f 0.6 10.6 f 0.6 10.6 f 0.7
The ratios k l / k 2and k 3 / k 2for the rate constants for chlorine atom [C2H6]/(ViBr]is shown in Figure 1 for competitions at three reactions with H2S or C& vs. CH2=CHBr are obtained from temperatures of 232,295, and 359 f 4 K. The reciprocal yields the slopes of the reciprocal yields of CH2=CH38Cl vs. the subare graphed with the assumption that 5% of the 38Clatoms are strate rataios of [H,S]/[ViBr] and [C2H6]/[ViBr], respectively. removed by hot reactions with CClF3 so that only 95% of the total The rate constant ratio k l / k 3is then determined indirectly from 38Clis available for reaction thermally;1,2the ratios of the slopes the relative efficiencies of HIS and C2H6 in diminishing the obare essentially independent of any error in the value of this nuserved yields of CH2=CH38C1 from CH2=CHBr. merator. The yields of CH2=CH38CI from CH2=CHBr vary Sample bulbs have been filled by standard vacuum line techwith temperature because they depend upon the competition niques with small ( l ) are negligible but provide no information about the rate for formation of HCl(u=O).8 The yield The progressive diminution in CH2=CH38C1 yields observed with increasing concentration ratios of [H2S]/ [ViBr] or for HCl(u=O) is certainly less than for HCI(u=l) because of the
J . Phys. Chem. 1986, 90, 1990-1992
1990
TABLE 11: Experimental Measurements of the Reaction Rate Constant for Chlorine Atoms with H2S temp, species press., 10"kl, cm3 K monitored methodo Torr molecule-I s-I ref 298 296 296 211-353 232-359
HCI C1 HZS CI CH2=CH3*CI
LP-CL DF-RF DF-MS FP-RF RGC
10 1-1.7 1 25-100 4000
7.3 f 0.9' 4.00 f 0.08 5.1 f 0.7 6.29 f 0.46 10.5 f 0.4
8 9 10 11 this
work LP, laser photolysis: CL, chemiluminescence: DF, discharge flow; RF, resonance fluorescence: MS, mass spectrometry; FP, flash photolysis; RGC, radiochemical gas chromatography. Rate Constant for formation of HCI(v=l). I?
intense laser action possible with this reaction, and the ratio (v=O)/(u=l) has been estimated as 0.6 from the work of Dill and Heydtmann.I5 If the latter estimate were correct, then the overall rate for kl from the data of Nesbitt and Leone would be (1.6) X (7.3 f 0.9) X lo-]', or (11.7 f 1.4) X lo-" cm3 molecule-l s d . The latter value is in agreement within the respective error limits of our measurement in Table I. The differences among the various values for k l in Table I1 appear to be beyond the respective error limits of the experiments, which in turn represent widely different experimental conditions and techniques. These radiochemical experiments are in a pressure ~~
~~
(15) Dill, B.: Heydtmann, H. Chem. Phys. 1978, 35, 161
regime at least 40 times higher than any of the other experiments, and our higher values may be indicative of a pressure dependence for k , . The only previous attempt to detect any change in k , with pressure used argon as the bulk gas and a pressure range of 25-100 Torr, well below that of our measurements." A possible mechanistic source of such variations can be that the reaction of C1 with H2S is not a direct abstraction of H but instead involves a more complex attack on the sulfur atom with the formation of an H2SCl intermediate, as suggested by Leone and colleagues.7*8 The absence of any observed activation energy in our experiments or in those of Nava et a1.I' does not provide much additional information about the mechanism of the reaction because both direct abstraction and complex formation can be made to occur with little temperature dependence, e.g., the abstraction of H from C,H, shown in Table I. An inconsistency also appears to exist among the various rate constant measurements at pressures of 100 Torr and less. In any case, the possibilities of a pressure dependence in the measured value for k , or of a complex intermediate H2SC1 surviving long enough to make collisions at 1-atm pressure introduce substantial uncertainty as to the appropriate value for k , to be used in atmospheric calculations or modeling. Our own experiments can presumably be modified to determine the value for k , with a mixture of N 2 / 0 2at 1-atm pressure as the moderator gas although some molecule such as CC1F3 must still be present as the source molecule for the 37Cl(n,y)38C1nuclear reaction.
Acknowledgment. This research was supported by the Office of Chemical Sciences, Division of Basic Sciences, U. S. Department of Energy, Contract No. AT-DE-ER76-70126.
A Novel Investigation of Vapor-Phase Charge-Transfer Complexes of Halogens with n Donors by Electron Energy Loss Spectroscopyt P. Vishnu Kamath, M. S. Hegde, and C. N. R. Rao* Solid State and Structural Chemistry Unit, Indian Institute of Science, Bangalore-560 012, India (Received: December 3, 1985; In Final Form: February 19, 1986)
Complexes of I2 with diethyl ether and triethylamine and of Br, with diethyl ether have been investigated in the vapor phase for the first time by employing electron energy loss spectroscopy. Besides the CT bands, blue-shifted vacuum-UV bands of the halogens have been assigned; the amine-I, system appears to exhibit two CT bands, associated with two different excited states of the complex.
Introduction Interactions between electron donors and acceptors are generally characterized by the occurrence of a charge-transfer (CT) tran~ition.'-~Besides the C T bands, in the case of halogen acceptors such as 12, a blue shift of the visible absorption band (520 nm) is ~ b s e r v e d . ~ The - ~ charge transfer and the blue-shifted visible I, bands have been investigated widely in many complexes of 1, in the solution p h a ~ e ,and ~ . ~the spectra of a few of the complexes have been characterized in the vapor We considered it most valuable to investigate the electronic spectra of some of the halogen complexes in the vapor phase by employing electron energy loss spectroscopys-I0 (EELS) since this technique would permit a study of the entire region from vacuum ultraviolet to visible. This would enable us to examine whether the vacuum UV transition of I2 (1 80 nm) also undergoes a blue shift on complexation. The interest in obtaining information on the vacuum UV transition of complexed I2 is due to the close rela'Contribution No. 342 from the Solid State and Structural Chemistry Unit.
tionship between the vacuum UV and the visible absorption bands.j The 180-nm band of iodine in complexes has not been studied hitherto probably because of the difficulty in recording optical spectra in the vacuum UV region. Another interest in the vacuum UV region is the possibility of observing an additional C T band (1) Mulliken, R. S.J. Am. Chem. SOC.1952, 74, 8 11. (2) Mulliken, R. S. J . Chim. Phys. 1963, 20. (3) Mulliken, R. S.; Person, W. B. Molecular Complexes: Wiley-Interscience: New York, 1969; p 142. (4) Breigleb, G. Electron-Donator-Acceptor Komplexe: Springer-Verlag: Berlin, 1961. ( 5 ) Rao, C. N. R.: Bhat, S. N.: Dwivedi, P. C. Appl. Spectrosc. Rev. 1971, 5, 1. (6) Rao, C. N. R.; Chaturvedi, G. C.; Bhat, S. N. J . Mol. Spectrosc. 1970, 33, 554. (7) Tames, M. In Molecular Complexes, Vol. 1, Foster, R. Ed., Elek Science: London, 1973; p 49. (8) Kupperman, A,; Flicker, W. M.; Mosher, 0. Chem. Rev. 1979, 79,77. (9) Rao, C. N. R.; Srinivasan, A,; Jagannathan, K.; Hegde, M. S. J . Sei. Ind. Res. (India) 1980, 39, 212. (10) Celotta, R. J.; Huebner, R. H. Electron Spectroscopy, Theory, Techniques and Applications, Brundle, C. R., Baker, A. D., Ed.; Academic Press: New York, 1979; p 41.
0022-3654/86/2090- 1990$01.SO10 0 1986 American Chemical Society
