130
J. Phys. Chem. 1996, 100, 130-137
Kinetic Studies of the Reaction between NO3 and OClO at T ) 300 K and P ) 2-8 Torr Andrew A. Boyd,† George Marston,‡ and Richard P. Wayne* Physical Chemistry Laboratory, South Parks Road, Oxford OX1 3QZ, U.K. ReceiVed: April 7, 1995; In Final Form: September 25, 1995X
The kinetics of the reaction between the radicals NO3 and OClO have been investigated at low pressure and room temperature using three laboratory techniques: discharge stopped-flow, slow-flow, and discharge-flow. In all types of experiment, strong evidence was found to suggest that reaction between the two species occurred on the walls of the reaction cell and led to the release of ClO into the gas phase. No evidence was found for the occurrence of the homogeneous reaction NO3 + OClO f products, and an upper limit for the rate coefficient for this process of 6 × 10-17 cm3 molecule-1 s-1 was determined. These experimental results are in disagreement with those of a previous study and suggest that this gas-phase reaction is unlikely to be of atmospheric importance even in the regions of the nighttime stratosphere where both species coexist.
Introduction The gas-phase reaction between NO3 and OClO
NO3 + OClO f products
(1)
is a potentially important process in Earth’s upper atmosphere.1 Both of these species have been identified as products of reactions known to be of stratospheric importance.2,3 NO3 is formed in the sequence of reactions
NO + O3 f NO2 + O2
(2)
NO2 + O3 f NO3 + O2
(3)
while OClO is generated Via the reactions
Cl + O3 f ClO + O2
(4)
Br + O3 f BrO + O2
(5)
ClO + BrO f OClO + Br
(6a)
Reactions 4 and 5 form part of a catalytic cycle for ozone destruction which is completed by another channel of reaction 6,
ClO + BrO f ClOO + Br
(6b)
followed by thermal decomposition of ClOO,
ClOO + M f Cl + O2 + M
(7)
Both NO3 and OClO have strong and characteristic electronic absorption spectra, which make them relatively easy to detect in small concentrations and to identify both in the nighttime atmosphere and in the laboratory. Consequently, temporally resolved measurements of NO3 concentration-altitude profiles have been used as evidence of the behavior of NOx and O3 above altitudes of 20 km, while polar OClO column abundances allow †Present address: Laboratoire de Photophysique et Photochimie Mole´culaire, Universite´ de Bordeaux I, 351 Cours de la Libe´ration, F-33405 Talence Cedex, France. ‡Present address: Department of Chemistry, University of Reading, Whiteknights, P.O. Box 224, Reading RG6 6AD, U.K. X Abstract published in AdVance ACS Abstracts, December 1, 1995.
0022-3654/96/20100-0130$12.00/0
inferences to be made about enhanced ozone destruction lower in the stratosphere that results from the ClOx/BrOx catalytic cycle.4 The generally good agreement between field measurements and modeling of the variations in stratospheric NOx and ClOx abundances has led to the belief that the main features of the chemistry and dynamics relevant to polar ozone depletion are relatively well understood. The only important nonphotolytic loss process for NO3 has been assumed to be indirect, through the formation of N2O5 and the subsequent heterogeneous removal of this species on polar stratospheric clouds (PSCs). OClO should be a short-term chlorine reservoir species which is stable at night in the polar spring. However, attempts to refine the modeling of polar ozone measurements have led to the conclusion that the existence of long-term chlorine reservoir species other than HCl and ClONO2 cannot yet be excluded.5 Suggested candidates include the higher chlorine oxides6 such as Cl2O6, which may be formed from OClO itself but which is more photolytically stable than its precursor.
OClO + O3 f sym-ClO3 + O2
(8)
ClO3 + ClO3 + M f Cl2O6 + M
(9)
Whether sym-ClO3 is a product of the reaction between OClO and O3 is not clear, and the rate coefficient for the overall process is known to be small.7 However, the reaction might still be significant under nighttime stratospheric conditions. The importance of NO3 and OClO to our understanding of stratospheric chemistry obviously makes it essential to characterize any as-yet unconsidered reactions which may affect their nighttime budgets and/or implicate these molecules in ozone depletion. This paper concerns laboratory studies of the potentially significant direct reaction between the two species,
NO3 + OClO f products
(1)
Previous studies in this laboratory have indicated that the reaction is slow. Discharge-flow experiments8 yielded an upper limit for the rate constant of 10-15 cm3 molecule-1 s-1, while a preliminary discharge stopped-flow study9 suggested a value of 2 × 10-16 cm3 molecule-1 s-1, assuming the reaction proceeded through the channel
NO3 + OClO f ClO + NO2 + O2 © 1996 American Chemical Society
(1a)
Kinetic Studies of the Reaction between NO3 and OClO
J. Phys. Chem., Vol. 100, No. 1, 1996 131
It was later concluded10 that, under stratospheric conditions, the homogeneous rate constant was probably even smaller than this value. Friedl et al.1 have since reported results obtained from studies of the NO3/OClO reaction system using a different technique. They utilized a slow-flow apparatus incorporating photodiodearray (PDA) and FTIR detection to monitor integrated UV and IR absorbances along the length of their cell, and they followed both reactant loss and product formation. Using a numerical simulation of a reaction scheme identical to that employed here, they modeled measured OClO and NO3 absorbances along with the absorbance of ClO, which they detected as a product. Their results appeared to confirm our assumption that reaction 1a is the only important channel at 298 K, but their derived value for k1 was 2 × 10-14 cm3 molecule-1 s-1, more than 1 order of magnitude greater than any of our previous estimates.8-10 An important result from the work of Friedl et al.1 was the discovery of unexpected IR and UV absorption features observed at temperatures below 250 K, which they attributed to the formation of the adduct O2ClONO2 (“chloryl nitrate”, ClNO5).
OClO + NO3 + M f O2ClONO2 + M
(1b)
Using a numerical model to describe their experimental results, Friedl et al. deduced that channel 1b becomes increasingly important at low temperatures and indeed dominates over channel 1a at stratospheric temperatures and pressures of a few Torrs. They therefore concluded that this adduct-forming channel could well be of significance in the “collar” region11 of the polar vortex, where processed and unprocessed polar air meet and hence where both reactants are likely to be present in the same nighttime air mass. The identification of the chloryl nitrate adduct by Friedl et al. appears indisputable, but their kinetic results seem to us to be open to further examination, as will be shown here. Furthermore, since they then used their measured values of k1b to estimate the bond strength in O2ClsONO2, and hence to estimate the diurnal variation and column abundance of ClNO5 in the collar region, the whole idea of ClNO5 being of any importance in the stratosphere is questionable. The inconsistency between their room temperature rate coefficient for reaction 1 and that presented here therefore becomes important. Three complementary techniques have been used to further study the OClO/NO3 reaction system, and the results from these discharge stopped-flow, slow-flow (similar to that of Friedl et al.), and conventional discharge-flow experiments are presented in turn. Experimental Section Materials. Chlorine dioxide was synthesized and purified in a Pyrex vacuum manifold, being produced at room temperature by the established method12 of passing molecular chlorine (BDH, >99.97%) through sodium chlorite (Aldrich). The yield of OClO was found to be greater if the NaClO2 was moistened and mixed with glass beads before packing into a U-tube, the beads ensuring a better exposure to the flowing Cl2. The emerging OClO was trapped out as a red solid at 77 K and then warmed to 195 K while pumping off any unreacted chlorine released in the process. Further purification was achieved8 by allowing the OClO to evaporate slowly and to recondense into another cold finger held at 77 K. Any crystalline residue of ice and orange solid (possibly an OClO clathrate) in the original finger was pumped away. Mass spectrometry was used to confirm the absence of any detectable higher chlorine oxide or
Cl2 impurities in the material prepared. The batch was stored at 77 K until required and then carefully allowed to warm until not more than 30 Torr of OClO vapor had been transferred into a bulb protected against explosion. The exposure of OClO to light was avoided at all times in view of its rapid photolysis in both the solid and gas phase.13 Careful dilution was carried out in the bulb to yield a mixture of 3-5% OClO in He. Such mixtures remained stable for several hours, during which time OClO could be added through a capillary flow meter in small but accurate amounts to the kinetic apparatus of interest. All three types of experiment used a buffer of He (B.O.C. Grade A, >99.9%), which entered the absorption cell or flow tube Via an oxygen-removing column (Oxisorb, Messer Griesheim). All other materials now to be described were used as supplied by the manufacturers, their flow rates being measured and controlled by the use of appropriate calibrated ball flow meters (Jencon, RS series) and needle valves (Nupro, S series). A fraction of the He flow was diverted through a nitric acid bubbler containing a 2:1 v/v mixture of 98% H2SO4 (BDH) and 70% HNO3 (BDH). The anhydrous nitric acid vapor subsequently formed nitrate radicals upon mixing with fluorine atoms, these having been generated upstream by passing a separate flow of 5% F2 in He (BOC) through a 2450 MHz microwave discharge.8 Discharge Stopped-Flow (DSF) Experiments. The apparatus and detection system used for the DSF experiments are shown schematically in parts a and b of Figure 1, respectively. The DSF technique for studying NO3 kinetics has already been reported in detail;14 only the principles and any alterations implemented for studying reaction 1 will be described here. The Pyrex reaction cell (22 cm long, 5.5 cm i.d.) was either internally coated with halocarbon wax or fitted with a PTFE sleeve in attempts to minimize surface reactions. A detection system incorporating a dual-boxcar analyzer was used to measure the integrated absorption signal of NO3 at λ ) 662 nm; the analysis beam was provided by a quartz-halogen lamp and interference filter arrangement and passed 12 times through the cell using White optics. Electronically controlled, custom-built Pyrex solenoid valves allowed the isolation of coreactant, at known initial concentrations, in the cell. Activation of these valves simultaneously triggered the recording of the decay in the NO3 absorption signal in real time over reaction times of several seconds or more. Several such absorption-time profiles recorded upon stopping the flow could be coadded until a suitable signal-to-noise ratio was achieved. The absorptions were subsequently converted to NO3 concentrations averaged through the cell ([NO3]av,t) at time t. NO3 concentrations were determined using the NO titration technique15 and were compared directly with measured absorbances. These calibrations yielded an effective σ662(NO3) in the range (1.1-1.5) × 10-17 cm2 molecule-1, assuming the nominal optical pathlength of 264 cm. The low values observed for the cross section3 were a reflection of the finite width of the filter (3.0 nm fwhm). Variability in the measured cross sections was caused by the dependence of the transmission profile of the filter on the angle between the filter and the incident radiation; for a fixed optical setup, the cross sections were reproducible within error. With the precursor flow rates shown in Figure 1a, typical calculated cell concentrations of NO3 averaged through the reaction cell before stopping the flow, [NO3]av,0, were in the range (1.52.5) × 1013 molecule cm-3. Slow-Flow (SF) Experiments. The principle of the slowflow experiments performed was very similar to that of the titration reaction of NO3 with NO used to measure the effective absorption cross section of NO3 at 662 nm for the DSF optical
132 J. Phys. Chem., Vol. 100, No. 1, 1996
Boyd et al.
Figure 1. Schematic diagram of (a, top) the apparatus and (b, bottom) the detection system used for the discharge stopped-flow experiments.14
setup.15,16 Most of the slow-flow experiments involved the same gas-handling system and cell as described previously for the DSF experiments but with a different optical detection system, the dual-beam setup with boxcar integrator being replaced by a photodiode array (PDA, Jobin-Yvon: 1024 diodes; diffraction grating with 200 grooves mm-1). The principal benefit of using a PDA was that the simultaneous monitoring of more than one absorbing species was possible both in the UV and the visible spectral regions. The Pyrex windows on the 5.5 cm diameter cell were replaced with quartz equivalents. The light source now consisted of a deuterium discharge lamp (Hamamatsu L1905 with stabilized power supply) whose output could, if necessary, be coupled to that of a quartz-halogen lamp, the D2lamp filament having a central iris through which collimated light could be passed from behind. After making two passes
through the absorption cell, the emerging beam was focused onto the entrance of a UV-transmitting fiber-optic cable, from where it directly entered the PDA. The concentrations of NO3 and OClO measured using the PDA system were calculated Via determinations of the effective absorption cross sections at chosen wavelengths. The electronic absorption spectrum of OClO was recorded, and its overall shape was found to agree very well with that measured elsewhere,17 only some fine structure at the shoulders of those peaks centered between λ ) 398 and 360 nm not being reproduced with the resolution of the PDA (1 diode being equivalent to ca. 0.7 nm in this region of the spectrum). The differential absorbance between the peak at λ ) 351 nm and the adjacent trough at λ ) 354 nm was used for calibration. Flows of OClO in He were added at measured rates to the cell, and the increase in the
Kinetic Studies of the Reaction between NO3 and OClO
J. Phys. Chem., Vol. 100, No. 1, 1996 133
absorbance through the cell was found for each addition (sampling time 10 s, single sample). A plot of the measured differential absorbance against [OClO] was linear and had no intercept, the slope being used to calculate an effective differential cross section of (8.2 ( 1.0) × 10-18 cm2 molecule-1 for OClO at 300 K (assuming an optical path length of 44.8 cm). The simplicity of this calibration procedure meant that the OClO cross section could be readily checked before each series of experiments. The mole fraction of the mixture of OClO in He was checked regularly during a series of experiments, by allowing say 10 Torr of the mixture from the bulb into a small passivated absorption cell and measuring the corresponding absorbance at λ ) 351 nm using a high-resolution UV/vis spectrometer (Perkin-Elmer, Lambda 5). For the detection of NO3, the band at λ ) 623 nm was used in preference to the slightly stronger band at λ ) 662 nm because of complicating second-order OClO peaks found to be present in the wavelength region 630-800 nm. The NO titration method15 was again used to estimate the effective absorption cross section for NO3 at this wavelength. The value of 8.0 × 10-18 cm2 molecule-1 obtained, compared with a recommended literature value3 of 1.4 × 10-17 cm2 molecule-1 at 298 K, reflects the convolution over the acceptance band of the PDA spectrometer.18 Discharge-Flow (DF) Experiments. A conventional discharge-flow apparatus was used to monitor the loss of NO3 as a function of time in the presence of OClO in a series of known excess concentrations. The DF arrangement has been described in detail elsewhere.19 The uncoated Pyrex flow tube used was 103 cm long and of 3.7 cm i.d. NO3 was generated at the rear of the flow tube, while OClO was admitted through a sliding injector. A recently introduced ballasted pumping setup20 allowed stable flow conditions to be maintained for linear flow velocities as low as 0.7 ms-1, and thus contact times of up to 1.2 s could be employed. NO3 concentrations were measured by absorption at λ ) 662 nm (the absorption cross section at this peak having again been calibrated by NO titration). The dual-beam multipath optical detection system was essentially the same as that already described for the discharge stoppedflow apparatus except that the absorbance was measured at a given point in flow rather than being integrated through the length of the reaction tube. The absorption path length in these experiments was 240 cm. Results and Their Interpretation Discharge Stopped-Flow Experiments. Figure 2a is an example of a recorded profile of the decay of NO3 alone in the cell at 300 K and for a typical total cell pressure of 8 Torr. There is an initial drop in the NO3 signal caused by the presence of a small amount of NO2 that reacts with NO3.
NO3 + NO2 + He f N2O5 + He
(10)
After this drop, the remainder of the decay can be attributed almost exclusively to loss on the cell surface,
NO3 + wall f products
Figure 2. NO3 decays in the discharge stopped-flow experiments: (a) [OClO] ) 0; (b) [OClO] ) 5.3 × 1012 molecule cm-3.
TABLE 1: Reaction Scheme Used for Kinetic Modeling of the NO3/OClO Reaction System process
k (300 K, 6 Torr He)a
ref
reaction number
Reactions Used for Friedl et al.1 OClO + NO3 f ClO + NO2 + O2 k1a (varied) ClO + NO3 f OClO + NO2 k15 ) 5.0 × 10-13 ClO + NO3 f Cl + O2 + NO2 β ) k15b/k15 (varied) Cl + NO3 f ClO + NO2 2.6 × 10-11 Cl + OClO f 2 ClO 5.8 × 10-11 ClO + OClO + M f Cl2O3 + M 1.1 × 10-14 NO2 + NO3 + M f N2O5 + M 2.2 × 10-13 ClO + ClO + M f Cl2O2 + M 4.1 × 10-15
(1a) 8, 17 (15a) (15b) 3 (16) 3 (17) 24 (18) 3 (10) 3 (19)
Additional Reactionsb Cl2O3 + M f ClO + OClO + M 100 s-1 ClO + ClO f products 1.2 × 10-14 Cl + NO2 + M f products 2.8 × 10-13 ClO + NO2 + M f ClONO2 + M 3.3 × 10-14 Cl + wall f products 5 s-1 ClO + wall f products 90% loss of the NO3 signal irrespective of temperature and the ratio of initial NO3 and OClO concentrations (varied between 5:1 and 1:5), while OClO concentrations dropped by much less. They attributed this observation to an initial gas-phase reaction between NO3 and OClO followed by a catalytic chain that destroys NO3 while conserving OClO. A more limited range of OClO concentrations was used in our slow-flow experiments performed using the DSF cell, although the data which Friedl et al. chose for kinetic modeling at 298 K did have values of [NO3]av similar to those listed in Table 2. The data sets that Friedl et al. elected to model were those with a cell pressure of 3 Torr and a residence time of 3-4 s. The reaction scheme they used was the suite of reactions listed in the top half of Table 1. They report a good match between experimental and simulated data for k1a and β values of 2 × 10-14 cm3 molecule-1 s-1 and 0.8, but no estimates are given of the errors in these derived parameters or of the consistency of the fit between several data sets. The value of k1a seems completely incompatible with our limits for a hypothetical gasphase process, and indeed with any of the entries in Table 3. However, in view of the results of the three types of experiment reported in the present paper, it seems difficult to imagine there not also being a heterogeneous effect involving OClO in the reaction system of Friedl et al. They used what was presumably an uncoated Pyrex cell, and although the diameter is 3 times that of our DSF cell, they perform experiments with 3 times the residence time and half of the cell pressure.
Boyd et al. TABLE 4: Comparison of Branching Ratios for Reaction 15 study
β
Biggs et al.8 Becker et al.22 Friedl et al.1 this work
0.8 ( 0.1 0.6 for all three types of experiment, and from the data tabulated in Tables 2 and 3 we give β ) 0.75 ( 0.16 (2σ); that is, the channel of reaction 15 leading to ClOO formation is more likely than regeneration of OClO. This conclusion is in line with the results of Biggs et al.8 and Friedl et al.,1 rather than those of Becker et al.22 The results of the various studies are drawn together in Table 4. Conclusions and Stratospheric Implications The kinetics of the reaction between NO3 and OClO has been studied by three different kinetic techniques (discharge stoppedflow, slow-flow, and conventional discharge-flow), with the results obtained strongly implying that the primary process is largely, if not exclusively, reaction between NO3 and OClO on the surface of the vessel. The need to implicate known secondary chemistry suggests that the products (ClO and NO2) desorb from the wall and react further with NO3. The room temperature results of Friedl et al.1 could not be repeated by us, even in slow-flow experiments very similar to the type they employed. There is thus an apparent inconsistency in the straightforward observations and not even in the interpretation. It may therefore be necessary to reconsider the results obtained by Friedl et al.1 at low temperature and especially to re-examine the conclusion that O2ClONO2 is a product of a gas-phase reaction. If the species were formed in
Kinetic Studies of the Reaction between NO3 and OClO the stratosphere at the rate Friedl et al. deduce from their kinetic results, then it could be a significant reservoir species in the collar region of the polar vortex. On the other hand, if the adduct is formed heterogeneously on the reactor wall before desorbing into the bulk, then the implications for the atmosphere are less clear. In the laboratory, the reactor wall activity itself may be governed by adsorbed HNO3 (one of the NO3 precursors) and traces of adsorbed H2O. Comparable chemistry could be envisaged on the surfaces of stratospheric aerosol. In that case, further quantitative studies of the surface reactivity of OClO may be of use not only in interpreting laboratory experiments but also in helping to understand heterogeneous processing in the stratosphere itself. Acknowledgment. The authors wish to thank Dr. C. E. Canosa-Mas, Ms. K. Hansen, Dr. P. S. Owen, and Mr. D. W. A. Stewart for their help during this investigation. R.P.W. and G.M. wish to thank SERC (NERC) and the CEC for grants to support some aspects of this work. References and Notes (1) Friedl, R. R; Sander, S. P.; Yung, Y. L. J. Phys. Chem. 1992, 96, 7490. (2) Wayne, R. P.; Barnes, I.; Biggs, P.; Burrows, J. P.; Canosa-Mas, C. E.; Hjorth, J.; LeBras, G.; Moortgat, G. K.; Perner, D.; Poulet, G.; Restelli, G.; Sidebottom, H. Atmos. EnViron. 1991, 25A, 1. (3) DeMore, W. B.; Golden, D. M.; Hampson, R. F.; Howard, C. J.; Kolb, C. E.; Kurylo, M. J.; Molina, M. J.; Ravishankara, A. R.; Sander, S. P. Chemical Kinetics and Photochemical Data for Use in Stratospheric Modeling: EValuation Number 11; JPL Publication 94-26, NASA Jet Propulsion Laboratory: Pasadena, CA, 1994. (4) Wayne, R. P. Chemistry of Atmospheres, 2nd ed.; OUP: Oxford, 1991. (5) Anderson, J. G.; Brune, W. H.; Lloyd, S. A.; Toohey, D. W.; Sander, S. P.; Starr, W. L.; Loewenstein, M.; Podolske, J. R. J. Geophys. Res. 1989, 94, 11480. (6) Friedl, R. R.; Sander, S. P. J. Phys. Chem. 1989, 93, 4756. Sander, S. P.; Friedl, R. R. J. Phys. Chem. 1989, 93, 4764.
J. Phys. Chem., Vol. 100, No. 1, 1996 137 (7) Wongdontri-Stuper, W; Jayanty, R. K. M.; Simonaitis, R.; Heicklen, J. J. Photochem. 1979, 10, 163. (8) Biggs, P.; Harwood, M. H.; Parr, A. D.; Wayne, R. P. J. Phys. Chem. 1991, 95, 7746. (9) Boyd, A. A.; Stewart, D. W. A.; Wayne, R. P. Stopped-Flow Studies of the Reactions of NO3 with OClO and O3; CHEMRAWN VIIth World Conference on The Chemistry of the Atmosphere: Its Impact on Global Change, Baltimore, MD, 2-6 Dec, 1991. (10) Boyd, A. A.; Stewart, D. W. A.; Wayne, R. P. Further StoppedFlow Studies of the Reactions of NO3 with OClO and O3, XIIth International Symposium on Gas Kinetics, Reading, U.K., 19-24 July, 1992. (11) Toon, J. C.; Farmer, C. B.; Lowes, L. L.; Schaper, P. W.; Blavier, J.; Norton, R. H. J. Geophys. Res. 1990, 94, 16571. (12) Derby, R. I.; Hutchinson, W. S. Inorg. Synth. 1953, 4, 152. (13) Hayman, G. D.; Cox, R. A. Chem. Phys. Lett. 1989, 55, 1. (14) Biggs, P.; Boyd, A. A.; Canosa-Mas, C. E.; Joseph, D. M.; Wayne, R. P. Meas. Sci. Technol. 1991, 2, 675. (15) Canosa-Mas, C. E.; Fowles, M.; Houghton, P. J.; Wayne, R. P. J. Chem. Soc., Faraday Trans. 1987, 83, 1465. (16) Boyd, A. A.; Canosa-Mas, C. E.; King, A. D.; Wayne, R. P.; Wilson, M. R. J. Chem. Soc., Faraday Trans. 1991, 87, 2913. (17) Wahner, A.; Tyndall, G. S.; Ravishankara, A. R. J. Phys. Chem. 1987, 91, 2734. (18) Mitchell, D. N.; Wayne, R. P.; Allen, P. J.; Harrison, R. P.; Twin, R. J. J. Chem. Soc., Faraday Trans. 2 1980, 76, 785. (19) Canosa-Mas, C. E.; Smith, S. J.; Toby, S.; Wayne, R. P. J. Chem. Soc., Faraday Trans. 2 1988, 84, 247. (20) Owen, P. S. D.Phil Thesis, University of Oxford, U.K., 1994. (21) Rattigan, O. M.; Jones, R. L.; Cox, R. A. Chem. Phys. Lett. 1994, 230, 121. (22) Becker, E.; Wille, U.; Rahman, M. M.; Schindler, R. N. Ber. Bunsen-Ges. Phys. Chem. 1991, 95, 1173. (23) Bagley, J. A.; Canosa-Mas, C. E.; Little, M. R.; Parr, A. D.; Smith, S. J.; Waygood, S. J.; Wayne, R. P. J. Chem. Soc., Faraday Trans. 1990, 86, 2109. (24) Burkholder, J. B.; Mauldin, R. L., III; Yokelson, R. J.; Solomon, S.; Ravishankara, A. R. J. Phys. Chem. 1993, 97, 7597. (25) Chance, E. M.; Curtis, A. R.; Jones, I. P.; Kirby, C. P. Report AERE-R 8775, Atomic Energy Research Establishment, Harwell, U.K., 1977.
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