11402
J. Phys. Chem. 1996, 100, 11402-11407
Low-Temperature Photochemistry of Submicrometer Nitric Acid and Ammonium Nitrate Layers Thomas G. Koch, Nicholas S. Holmes, Tristan B. Roddis, and John R. Sodeau* School of Chemical Sciences, UniVersity of East Anglia, Norwich, NR4 7TJ, England ReceiVed: February 7, 1996; In Final Form: April 25, 1996X
Reflection-absorption infrared spectroscopy has been employed in order to investigate the low-temperature photochemistry (90-140 K) of thin films of nitric acid and ammonium nitrate grown in Vacuo. Photolysis of amorphous nitric acid hydrate, the crystalline dihydrate (NAD) and trihydrate (NAT) at λ > 230 nm resulted in the formation of molecular nitric acid due to rapid protonation of the excited nitrate ion. Secondary photolysis of HONO2 produced NO2 and NO. If a neat film of molecular, anhydrous nitric acid was irradiated, nitrate and nitronium ions were observed. In contrast, ammonium nitrate photolysis at 140 K did not result in a proton transfer to produce NH3 and HONO2 but in the formation of the peroxynitrite ion (ONOO-) as a precursor for NO2-. Molecular dinitrogen tetraoxide and nitrous oxide were also detected in the film. Mechanistic details and possible implications for the chemistry of the polar atmosphere are discussed.
Introduction The nitrate ion exhibits a characteristic UV absorption spectrum consisting of a strong band centered at 200 nm due to an allowed π-π* transition (max ) 9.9 × 103 mol-1 cm-1) and a considerably weaker absorption at 300 nm due to a forbidden n-π* transition (max ) 7.3 mol-1 cm-1).1 Therefore the photolysis of inorganic nitrates has been studied extensively in aqueous and solid media employing conventional analytical methods2-8 as well as a variety of spectroscopic techniques, such as electron spin resonance9,10 and infrared spectroscopy.11-13 Photodecomposition to produce peroxynitrite ions, nitrite ions, and nitrogen di- or monoxide has been suggested to occur via channels 1a-d following photon absorption.
NO3- + hν f OONO-
(1a)
f NO2- + O
(1b)
f NO2 + O-
(1c)
f NO + O2-
(1d)
Nitric oxide formation via channel (1d) has recently been observed in laser ablation studies of crystalline sodium nitrate.14 Infrared evidence suggests further that nitrite ions may be produced not only directly via (1b) but also via peroxynitrite as an intermediate.11 The peroxynitrite ion is known to absorb light near 325 and 375 nm and may reisomerize to the nitrate ion or dissociate to form NO2-.15,16
OONO-+ hν f NO3f NO2- + O
(2a) (2b)
In a diffuse reflectance infrared study, Vogt and FinnlaysonPitts13 have investigated the photochemistry of amorphous surface nitrate as formed by heterogeneous atmospheric reactions of nitrogen oxides on seasalt aerosols. Contrasting previous work on crystalline nitrates, they reported the formation of NO2 X
Abstract published in AdVance ACS Abstracts, June 1, 1996.
S0022-3654(96)00368-1 CCC: $12.00
rather than NO2- or ONOO- as the major photolysis channel. NO2 radicals have also been detected by ESR spectroscopy upon the irradiation of ammonium nitrate10 and frozen nitric acid,9 both of which may also be found in the atmosphere as part of aerosols17 or polar stratospheric clouds.18 The photodecomposition of ammonium nitrate is thought to be accompanied by the formation of nitrous oxide, although a mechanism has not been reported.4 In this paper, we describe a series of experiments, in which we have investigated the low-temperature photochemistry of thin films of nitric acid and ammonium nitrate. We observe a distinct difference in the photolytic decomposition of the nitrate ion in the two environments which is largely determined by the reactivity of the two counterions, H3O+ and NH4+, respectively. Experimental Section The experimental arrangement used in this study has been detailed elsewhere.19 Reflection-absorption infrared spectroscopy was employed for the analysis of thin substrate films deposited in Vacuo. The infrared beam emitted from a Digilab FTS-60 spectrometer was reflected off a temperature-controlled gold substrate at a high angle of incidence to the surface normal (ca. 75°) and detected by an external liquid nitrogen cooled MCT detector. RAIR spectra were recorded at 4 cm-1 resolution from co-addition of 256 interferograms. Based on infrared intensities and exposure (1 × 10-6 mbar s-1 ) 1 langmuir ≈ 1 monolayer s-1), the layers were estimated to be of the order of 100 nm thick. Background pressures in the differentially pumped vacuum chamber typically ranged from 1 × 10-9 to 5 × 10-9 mbar. On the basis of data of Hanson and Mauersberger,20 thin layers of amorphous nitric acid hydrate, the crystalline nitric acid dihydrate (NAD), and trihydrate (NAT) were grown by depositing the vapor of concentrated nitric acid solutions (at 1 × 10-6-1 × 10-5 mbar for 1-5 min) onto the gold substrate at 160-190 K. Anhydrous nitric acid was synthesized by vacuum distillation of a sodium nitrate (Fisons, stated purity 98%) and concentrated sulfuric acid (Fisons, 98%) mixture according to a method by Johnston.21 The vapor was deposited at 90 K (1 × 10-7-1 × 10-6 mbar for 1-2 min) and briefly annealed to 140 K in order to offset any annealing effects due © 1996 American Chemical Society
Photochemistry of HNO3 and NH4NO3 Thin Layers
J. Phys. Chem., Vol. 100, No. 27, 1996 11403 TABLE 1: Vibrational Assignment of Nitric Acid Hydrates normal mode
NAX/cm-1
NAD/cm-1
NAT/cm-1
ν1,3(H2O)
3400 vs
ν4(H3O+) ν3(NO3-)
1750 m 1450 m 1350 m
3490 vs 3220 m 1750 vw 1460 m 1270 m 1180 vs 1030 w 810 s
3430 vs 3250 vs 1790 vw 1390 m
ν2(H3O+) ν1(NO3-) ν2(NO3-)
Figure 1. RAIR spectra of amorphous nitric acid hydrate (A), crystalline nitric acid dihydrate (B), and crystalline nitric acid trihydrate (C) on gold foil at 90 K.
to irradiation. Ammonium nitrate films were prepared in situ by co-depositing the vapor of a concentrated nitric acid solution with ammonia (BDH, 99%) at 180-200 K (total pressure, 1 × 10-6-1 × 10-5 mbar for 1-5 min). Ammonia, which does not stick to the gold surface above 110 K under our experimental conditions, was used in excess to prevent the condensation of nitric acid hydrates. Ammonium nitrate films were briefly annealed to 270 K to remove any residual water. All substrates were routinely cooled to 90 K before irradiation with the continuous output of a 70 W xenon lamp (λ > 230 nm) which caused annealing to 120-140 K. For photolysis at λ > 300 nm a 5 mm thick Pyrex glass filter was inserted into the light beam. In the present experimental setup, quantitative measures of photochemical processes such as as quantum yields cannot be determined and therefore the results discussed below provide only mechanistic insights about nitrate photolysis in water-ice films. Results RAIR spectra of amorphous nitric acid hydrate (NAX), NAD, and NAT are shown in Figure 1, traces A, B, and C, respectively. Infrared absorptions due to the components H3O+, NO3-, and H2O are listed in Table 1. Although the peak positions of the crystalline hydrates, NAD and NAT, reported here agree well with data in the literature, the relative intensities differ considerably.22 In particular, the ν2(NO3-) nitrogen outof-plane and the ν2(H3O+) umbrella mode appear to be very strong by comparison whereas the ν3 (NO3-) asymmetric
1040 w 820 w
1130 m 820 m
stretching and ν4(H3O+) bending mode are very weak. Such a significant change in the relative band intensity on going from the transmission to the RAIRS experiment is usually indicative of an orientation effect since the underlying gold surface imposes the so-called “metal surface selection rule”:23 only vibrations with a dipole moment change perpendicular the surface can be detected because of the screening effect of the metal valence electrons. Both the ν2(NO3-) and ν2(H3O+) modes induce a dipole along their C3 axes whereas both the ν3(NO3-) and ν4 (H3O+) modes have a changing dipole perpendicular to that axis. The change observed in the relative intensities can therefore be explained if both the molecules preferentially order with their C3 axes perpendicular the metal surface. Such an alignment would cause the relative intensities of the two ν2 modes to increase, as their dipole moments align with the electric field. At the same time the ν3(NO3-) and ν4(H3O+) modes become infrared inactive in RAIRS. Evidently, this phenomenon was not observed for amorphous nitric acid hydrates due to random ordering. A more comprehensive FTIR study regarding the alignment of nitrate ions in/on ice surfaces is currently being carried out. The absorption difference spectra of amorphous nitric acid hydrate and crystalline NAD and NAT recorded after 5 h of photolysis at λ > 230 nm are shown in Figure 2, traces A, B, and C, respectively. Negative bands denote the loss of the precursor, whereas positive absorptions arise from the formation of photoproducts. In all three cases, molecular nitric acid was formed as evidenced by a new band near 1690 cm-1 due to the strong asymmetric NO2 stretch. In trace A, additional absorptions at 1310, 964, and 770 cm-1 were detected due to the symmetric NO2 stretch, N-O stretch, and ONO2 bending mode, respectively. Furthermore, weak bands due to the formation of monomeric NO2 (1610 cm-1) and NO (1860 cm-1) were observed in trace A. In the cases of the NAD and NAT, photolysis was accompanied by a pronounced loss of the sharp bands of the water of crystallization. Presumably, this does not arise from the actual photolysis of H2O, as water does not absorb light above 200 nm, but rather the destruction of the well-defined, crystalline H3O+(H2O)n)1,2NO3- complexes due to molecular HONO2 formation. The broad band observed in the 3400 cm-1 region is therefore likely to be the result of the formation of disordered water. However, ice forming on the liquid nitrogen cooled MCT detector after prolonged photolysis times may also contribute to this band. The photochemistry of molecular nitric acid was investigated in a separate experiment. Figure 3, trace A, shows the RAIR spectrum of a neat film of covalent HONO2 (3150, 1690, 1335, 964, and 772 cm-1). Trace B shows the absorption difference spectrum after photolysis. Characteristic absorptions due to NO2+ (2373 cm-1) and NO3- (1430, 1334, and 1040 cm-1) were readily identified although the bands of the nitrate ion partly overlap with the molecular HONO2 parent peaks near 1400 cm-1. The absorption at 1710 cm-1 was tentatively assigned to the NdO stretch of peroxynitrous acid by comparison with previous work.24-26
11404 J. Phys. Chem., Vol. 100, No. 27, 1996
Figure 2. RAIR difference spectra of amorphous nitric acid hydrate (A), crystalline nitric acid dihydrate (B), and crystalline nitric acid trihydrate (C) after 300 min photolysis at λ > 230 nm (T ) 120 K).
Koch et al.
Figure 4. RAIR spectra of ammonium nitrate as deposited at 190 K (A), annealed to 270 K (B), and recooled to 90 K (C).
TABLE 2: Vibrational Assignment of Ammonium Nitrate at Various Temperatures NH4+NO3-/cm-1 normal mode
at 190 K
ν3(NH4+) ν1(NH4+) ν2(NH4+) ν4(NH4+)a
3260 s 3081 m 1760 vw 1471 vs
3254 s 3087 m 1771 vw 1479 vs
ν3(NO3-)a
1390 m
1350 m, sh
ν1(NO3-) ν2(NO3-)
1046 vw 826 w
1051 vw 833 m
a
Figure 3. RAIR spectrum of anhydrous nitric acid as deposited at 90 K (A), and after 60 min photolysis at λ > 230 nm (B) (difference spectrum).
RAIR spectra of ammonium nitrate films as grown at 190 K, annealed to 270 K, and recooled to 90 K are shown in Figure 4, traces A, B, and C, respectively. The vibrational assignments are listed in Table 2. We found that relative intensities and splitting patterns, particularly in the 1300-1500 cm-1 spectral region, were extremely variable with temperature and deposition pressures as has been observed by other authors.27,28 These observations are not entirely surprising since the ν4(NH4+) and ν3(NO3-) modes have very similar frequencies and hence are prone to site-sensitive coupling. Optical effects, as we have previously observed in a RAIR study of thin films of dinitrogen tetraoxide and dinitrogen pentaoxide,29 may also contribute to
at 270 K
at 90 K 3240 s 3072 m 1775 vw 1492 vs 1462 vs 1397 m 1367 m 1320 m 1055 vw 833 m
The two modes may be coupled.
frequency shifts and intensity transfer. Furthermore, ammonium nitrate exists in at least six different phases although only the tetragonal, phase V is thought to be the stable between 75 and 257 K.28 Despite variable infrared spectra of the precursor we did not observe any inconsistencies concerning the photolytic decomposition pattern in the low-temperature photolysis experiments (90-140 K) reported here. Figure 5 shows a typical RAIR difference spectrum of ammonium nitrate after a 60 min photolysis with the unfiltered output of the xenon lamp at 140 K. A new set of bands at 1558, 962, and 790 cm-1 was observed, which was assigned to the NdO stretch of the trans conformers and the OO stretching and ONO bending modes of the cis conformer of ONOO-, respectively, in excellent agreement with recent work by Lo et al.12 The absorption at 1226 cm-1 was attributed to the asymmetric stretch of the nitrite ion being similar to its argon
Photochemistry of HNO3 and NH4NO3 Thin Layers
J. Phys. Chem., Vol. 100, No. 27, 1996 11405 Discussion Nitric Acid. The primary step in any photochemical process is the absorption of a photon. Walsh 30 was the first to suggest that if a planar AB3 molecule with 24 valence electrons, such as the nitrate ion, absorbs a photon it becomes pyramidal in its excited state.
(NO3-)planar + hν f (NO3-)*pyramidal
Figure 5. RAIR difference spectrum of ammonium nitrate after 60 min of photolysis at λ > 230 nm (T ) 140 K).
matrix value of 1244 cm-1. N2O (2235 cm-1) and N2O4 (1745, 1260, and 750 cm-1) were also readily identified. In addition, absorptions at 1313 and 1040 cm-1 were observed near the nitrate parent peak absorption. They are characteristic of the lower frequency component of the normally doubly degenerate asymmetric stretch and totally symmetric stretch of nitrate ions, respectively, formed in a different, less symmetric environment.13 The higher frequency component of the ν3(NO3-) mode near 1450 cm-1 is thought to overlap with the bands of the ammonium ion. After irradiation, films remained unchanged over time if kept in the dark. The normalized kinetic profiles of the photolysis products are illustrated in Figure 6. The steep initial slope and strongly curved profile observed for the formation of the peroxynitrite ion suggests that it serves as a precursor for the subsequent production of NO2-. The kinetics of N2O4 and N2O are necessarily more complicated as their production depends on the decomposition of two precursor ions. Photolysis experiments using a Pyrex filter only transmitting wavelengths above 300 nm were unsuccessful over a period of 8 h. It is still unclear whether this is due to the reduced photon flux and low maximum extinction coefficient of the n-π* transition (max ) 7.3 mol-1 cm-1) or to the fact that only excitation of the more energetic π-π* transition results in decomposition. Laser photolysis experiments are planned for the near future to investigate this issue further.
(1′)
The change of symmetry results in a decrease in the bonding character along the N-O bonds and an increase in bonding character in the O-O direction which can lead to intramolecular rearrangement or dissociation via channel 1a-d to produce ONOO-, NO2-, NO2, or NO, respectively. Upon UV irradiation of nitrate ions as part of amorphous or crystalline nitric acid hydrate films, we observe molecular HONO2 as the only initial product. This discrepancy can be most easily be explained by a rapid proton transfer from the oxonium ion to the highly reactive, excited nitrate ion preventing subsequent decomposition or isomerization.
H3O+•nH2O + (NO3-)* f HONO2 + (n+1)H2O (3) Numerous previous studies have reported low photolysis yields of nitrates in dilute, aqueous solutions below pH 7 because of quenching of the excited state by protonation.8 However, molecular nitric acid is highly unstable under these conditions and dissociates quickly to reform the precursor. In our lowtemperature experiments, reionization of HONO2 by the surrounding water molecules is largely prevented by the low activation energy available, making this process irreversible. The additional products, NO and NO2, are only observed after several hours of photolysis. This indicates that they are unlikely to be primary products but rather due to the secondary photolysis of molecular nitric acid, which itself absorbs light below 300 nm.31
HONO2 + hν f OH + NO2
(4)
NO2 + hν f NO + O
(5)
Reactions 4 and 5 are well established in nitric acid gas phase photochemistry. However, we do not observe the thermodynamically less stable isomer of nitric acid, peroxynitrous acid.
Figure 6. Normalized kinetic profiles of ONOO-, NO2-, N2O, and N2O4 produced upon ammonium nitrate photolysis.
11406 J. Phys. Chem., Vol. 100, No. 27, 1996
Koch et al.
HOONO has previously been observed upon HONO2 photolysis in inert low-temperature matrices (4-12 K). Such a product is thought to be formed from radical recombination.24-26
OH + NO2 + M f HOONO + M
(6)
In the present study, hydrogen bonding to water presumably prevents efficient recombination of OH and NO2 and the concentration of HOONO may simply be too low to be detected. In contrast to the above results on nitric acid hydrate photolysis, the irradiation of a neat film of anhydrous nitric acid leads to the formation of HOONO, NO2+, and NO3-. Thus, upon photolysis of neat molecular nitric acid, the OH radical formed via reaction 4 can react either with NO2 to produce peroxynitrous acid or with an adjacent HONO2 molecule to form H2O and NO3.
OH + HONO2 f H2O + NO3
(7)
The latter process is consistent with a previous ESR study which detected OH, NO2, and NO3 radicals upon photolysis of a frozen 6 M nitric acid solution.9 Reaction of NO2 and NO3 radicals can then produce N2O5 which rapidly ionizes to NO2+ and NO3at temperatures above 90 K.19
NO2 + NO3 f [N2O5]q f NO2+ + NO3-
(8)
This result dramatically highlights potential differences between atmospheric gas phase and heterogeneous photochemistry. Ammonium Nitrate. Ammonium nitrate displays a photochemistry distinctly different from nitric acid. We observe a number of decomposition products rather than the facile formation of NH3 + HONO2 which would be produced by a proton transfer. Presumably, this finding results from the lower acidity of the ammonium ion compared to the oxonium ion. In the experiments reported in this study peroxynitrous acid appears to be formed via (1a) as the only primary product. The kinetic profiles shown in Figure 6 suggest, further, that it is a likely precursor for the formation of nitrite.
NO3- + hν f OONO-
(1a)
OONO- + hν f NO2- + O
(2b)
Similar results have also been obtained by Plumb and Edwards upon irradiation of potassium nitrate crystals.11 They also reported that the steady state concentration of peroxynitrite is largely controlled by rephotoisomerization.
OONO- + hν f (NO3-)′
(2a′)
be consistent with the relatively long induction period observed in our study for the formation of dinitrogen tetroxide. As illustrated in Figure 6, N2O formation also commences after a time delay, which suggests that it may be formed by a similar mechanism. However, unlike N2O4 the formation of N2O requires decomposition of the nitrate ion as well as the ammonium ion. A likely process is the reaction of ammonium with the highly reactive O- atom to produce NH2 and H2O, which is exothermic by 849 kJ mol-1. The NH2 radical may then go on to react with nitrogen dioxide formed together with O- via (1c) and so produce nitrous oxide and water.32-34
NH4+ + O- f NH2 + H2O
(10)
NH2 + NO2 f N2O + H2O
(11)
Unfortunately, we are unable to observe NH2 in the thin-film infrared spectra, presumably because it is a weak absorber and the bands overlap with the ammonium precusor. Furthermore, by analogy to ammonia, which desorbs at 110 K, it is likely that a certain amount of NH2 escapes into the gas phase at 140 K. A low concentration of NH2 would, of course, also promote N2O4 formation. Conclusion The effect of heterogeneous photochemistry on atmospheric chemistry has not been included hitherto in models predicting chemical composition changes in the polar regions. The main reason for this omission is understandable because few photoprocesses have been studied in the laboratory under appropriate mimic conditions. Indeed, even the UV-visible spectra of relevant species adsorbed to water-ice or aerosol surfaces are not available in the literature. In this study we have presented some preliminary steps in a systematic exploration of photoinitiated chemistry, which may be pertinent to polar ozone research. We have demonstrated for the first time that in the case of ammonium nitrate the peroxynitrite ion is a likely precursor not only for nitrite ions but also for molecular NO2 and N2O. Ammonium nitrite exists in a thermodynamic equilibrium with ammonia and nitrous acid; therefore, HONO as well as NO2 and N2O may be released from aerosol surfaces under the higher temperatures prevailing in the troposphere. However, heterogeneous photochemistry with potential impact on the atmosphere involving nitrate ions is likely to be confined to neutral or alkaline substrates. In strongly acidic environments, as provided by type I polar stratospheric cloud particles or sulfuric acid acid containing aerosols, effective inorganic nitrate photolysis will be largely prevented due to proton transfer in the excited state.
The nitrate ion, regenerated via reaction 2a may possess a lattice orientation different from the precursor. Such a suggestion may explain the observed formation of “less symmetric” nitrate ions upon irradiation (1313 and 1040 cm-1). Vogt and Finnlayson-Pitts13 have studied similar, amorphous surface nitrate species which were found to decompose preferentially to produce NO2 radicals rather than OONO-.
Acknowledgment. This work has been carried out with support from the NERC LSAC Programme as well as the CEC STEP and Environment programmes. We also thank Dr. Andrew Horn for many useful discussions over the last few years.
(NO3-)′ + hν f NO2 + O-
(1) Blandamer, M. J.; Fox, M. F. Chem. ReV. 1970, 70, 59. (2) Doigan, P.; Davis, T. W. J. Phys. Chem. 1952, 56, 764. (3) Papee, H. M.; Petriconi, G. L. Nature 1964, 204, 142. (4) Zawidsky, T. W.; Petriconi, G. L.; Papee, H. M. Nature 1965, 207, 262. (5) Daniels, M.; Meyers, R. V.; Belardo, E. V. J. Phys. Chem. 1968, 72, 389. (6) Shuali, U.; Ottolenghi, M.; Rabani, J.; Yelin, Z. J. Phys. Chem. 1969, 73, 3445.
2NO2 + M f N2O4 + M
(1c′) (9)
At low temperatures, recombination of two NO2 molecules via reaction 9 inevitably leads to the formation of dinitrogen tetraoxide,29 which we observe. Such a mechanism would also
References and Notes
Photochemistry of HNO3 and NH4NO3 Thin Layers (7) Barat, F.; Gilles, L.; Hickel, B.; Sutton, J. J. Chem. Soc. A 1970, 1982. (8) Baylis, N. S.; Bucat, R. B. Aust. J. Chem. 1975, 28, 1865. (9) Hayon, E.; Sato, E. J. Chem. Phys. 1965, 43, 4314. (10) Pace, M. D. J. Phys. Chem. 1994, 98, 6251. (11) Plumb, R. C.; Edwards, J. O. J. Phys. Chem. 1992, 96, 3245. (12) Lo, W.-J.; Lee, Y.-P; Tsai, J.-H. M.; Tsai, H.-H.; Hamilton, T. P.; Harrison, J. G.; Beckman, J. S. J. Chem. Phys. 1995, 103, 4026 (13) Vogt, R.; Finlayson-Pitts, B. J. Phys. Chem. 1995, 99, 17269. (14) Bradley, R. A.; Lanzendorf, E.; McCarthy, M. I.; Orlando, T. M.; Hess, W. P. J. Phys. Chem. 1995, 99, 11715. (15) Krauss, M. Chem. Phys. Lett. 1994, 222, 514. (16) Lo, W.-J.; Lee, Y.-P; Tsai, J.-H. M.; Beckman, J. S. Chem. Phys. Lett. 1995, 242, 147. (17) Wall, S. M.; John, W.; Ondo, J. L. Atmos. EnViron. 1988, 22, 1649 (18) Turco, R. P.; Toon, O. B.; Hamill, P. J. Geophys. Res. 1989, 94, 16493. (19) Horn, A. B.; Koch, T.; McCoustra, M. R. S.; Chesters, M. A.; Sodeau, J. R. J. Phys. Chem. 1994, 98, 946. (20) Hanson, D.; Mauersberger, K. J. Phys. Chem. 1988, 92, 6167. (21) Johnston, H. S.; Cheng, S.-G.; Whitten, G. J. Phys. Chem. 1974, 78, 1.
J. Phys. Chem., Vol. 100, No. 27, 1996 11407 (22) Koehler, B. G.; Middlebrook, A. M.; Tolbert, M. A. J. Geophys. Res. 1992, 97(D8), 8065 (23) Pearce, H. A.; Sheppard, N. Surf. Sci. 1976, 59, 205. (24) Koch, T. G.; Sodeau, J. R. J. Phys. Chem. 1995, 99, 10824 (25) Chen, W.-J.; Lo, W.-J.; Cheng, B.-M.; Lee, Y.-P. J. Chem. Phys. 1992, 97, 7167. (26) Cheng, B-M.; Lee, J.-W.; Lee, Y. P. J. Phys. Chem. 1991, 95, 2814. (27) Keller, W. E.; Halford, R. S. J. Chem. Phys. 1949, 17, 26 (28) Shen, Z. X.; Kuok, M. H.; Tang, S. H. Spectrochim. Acta, 1993, 49A, 21 (29) Koch, T. G., Horn, A. B.; McCoustra, M. R. S.; Chesters, M. A.; Sodeau, J. R. J. Phys. Chem. 1995, 99, 8362. (30) Walsh, A. D. J. Chem. Soc. 1953, 2301. (31) Burkholder, J. B.; Talukdar, R. K.; Ravishankara, A. R.; Solomon, S. Geophys. Res. 1993, 98D, 22937. (32) Hack, W.; Schacke, H.; Schroter, M.; Wagner, H. G. EleVenth Symp. (Int.) Combust. 1979, 17, 505. (33) Kurasawa, H.; Leschaux, R. Chem. Phys. Lett. 1979, 66, 602. (34) Xiang, T. X.; Torres, L. M.; Guillory, W. A. J. Chem. Phys. 1985, 83, 1623.
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