J . Phys. Chem. 1991,95, 5815-5821
5815
Atmospheric Fate of CF2H2,CH3CF3,CHF2CF3,and CH3CFCi2: Rate Coefficients for Reactions with OH and UV Absorption Cross Sections of CH3CFCi2 Ranajit Talukdar, Abdelwahid Mellouki, Tomasz Gierczak3 James B. Burkholder, Stuart A. McKeen, and A. R. Ravishankara* NOAA, Aeronomy Laboratory, 325 Broadway, Boulder, Colorado 80303, and Cooperative Institute for Research in Environmental Sciences, University of Colorado, Boulder, Colorado 80309 (Received: December 6, 1990; In Final Form: March 5, 1991)
The absolute rates coefficients for reactions of hydroxyl radical with four halocarbons (CH2F2(HFC-32), CH3CF3(HFC-143a). CF3CHF2(HFC-I25), and CH3CFC12(HCFC-I4lb)), which are potential substitutes for CFCs, were measured by using the discharge flow-laser magnetic resonance and pulsed photolysis-laser-induced fluorescence techniques between 222 and 393 K. The UV absorption cross sections for HCFC-14lb between 190 and 230 nm were also measured as functions of temperature. These kinetics and cross-section data are compared with results from previous studies. These data were incorporated into a one-dimensional photochemical transport model to estimate the tropospheric and atmospheric lifetimes needed to assess the ozone depletion potential and the global warming potential of these substitutes. The atmospheric lifetime of CH3CFCI2 (HCFC-141b) is approximately 50% longer than the previously accepted value.
OH
Introduction
Chlorofluorocarbons (CFCs) are inert in Earth's troposphere and are transported into the stratosphere where photolysis releases active chlorine.' The active chlorine can catalytically deplete ozone2 in the stratosphere. Because they are inert, these compounds also accumulate in the atmosphere and contribute to the greenhouse effect.' Because of these two environmental concerns, the production of CFCs is restricted by an international agreement, the Montreal Protocol on substances that destroy stratospheric ozone.' Consequently, search for potential substitutes for chlorofluorocarbons are underway. The partially halogenated ethanes, hydrochlorofluorocarbons (HCFCs), and hydrofluorocarbons (HFCs), are currently prime replacement candidates because of their physical properties5 and the expected lower ozone depletion potential. The presence of an H atom in HCFCs and HFCs makes them reactive toward OH and have shorter lifetimes in the troposphere. Hence, a smaller fraction of the released compounds will be transported into the stratosphere and deplete ozone to a lesser extent than the CFCs. In this argument it is implicitly assumed that, once the degradation of the molecule is initiated in the troposphere, no chlorine-containing product that is stable enough to be transported to the stratosphere is produced. Of course, HFCs do not contain any chlorine and hence have a zero ozone depletion potential (ODP).However, they would act as greenhouse gases if retained in the atmosphere. The HFC and HCFCs are water insoluble and are unreactive with common tropospheric oxidizers such as ozone. Photolysis, as shown here, can contribute to their loss only in the stratosphere. Therefore, reactions with hydroxyl radicals (OH) are likely to be the main homogeneous gas-phase removal processes for the substitutes in the troposphere. We have measured the absolute rate coefficients for their reactions with O H and their UV absorption cross sections necessary to calculate atmospheric removal rates. Reactions with O(lD) can also play a role in the stratospheric removal of these compounds and are currently being studied in our laboratory. In previous papers we reported the rate coefficients6 for reactions with OH and the UV absorption cross sections' for five of the proposed substitutes, CF3CH2F(HFC134a), CF3CHCIF (HCFC-124), CF3CHC12 (HCFC-123), CH3CHF2 (HFC-l52a), and CH3CF2Cl(HCFC-142b). In this paper we present the results of a kinetic study of the reactions of OH with four other substitutes: A d d m all correspondenceto this author at NOAA. Also affiliated with the Department of Chemistry and Biochemistry, University of Colorado, Boulder, CO 80309. 'on leave from Department of Chemistry, Warsaw University, Zwirki i Wigury 101. 02-089. Warsaw, Poland.
0022-3654/91/2095-58 15$02.50/0
+ CF2H2 (HFC-32)
-
+
H20
+ CF2H;
kl
+ CH3CF3 (HFC-143a) H 2 0 + CH2CF3; k2 O H + CHF2CF3 (HFC-125) H20 + CF2CF3; k3 OH + CH3CFC12 (HCFC-14lb) H2O + CH2CFC12; OH
+
.-+
(1) (2) (3) k4
(4) The UV absorption cross sections for CH3CFCI2are also reported. The rate coefficients, and in particular their temperature dependencies, for the above four reactions had not been extensively studied until the recent announced use of these compounds as CCF replacement^.^ Since then there have been more investigations of reactions 1-4. The previous measurements of kl-k4 are listed in Table 111 (for CH2F2),Table V (for CH3CF3),Table VI1 (for CF3CHF2),and Table IX (for CH3CFCI2). The existing data have been There were no published values of k4when our investigation was initiated. Concurrent with our work, the investigations of Liu et aL9 and Brown et al.10were carried out. Our preliminary results on k4agreed with those of Liu et al., as reported in the AFEAS reportas However, we suspected the presence of a reactive impurity and hence reinvestigated k4,which is reported here. We have studied the above reactions between 222 and 393 K using two complementary methods, the pulsed photolysis and the flow tube techniques. The UV absorption cross sections of CH3CFC12 have been measured previously by Gillotay et al." at 298 K. The temperature (1) Rowland, F. S.;Molina, M. J. Rev. Geophys. Space Phys. 1975, 13, 1-36. (2) See for example: Brasseur, G.; Solomon, S. Aeronomy of rhe middlc
atmosphere; D. Reidel: Dordrccht, 1984. (3) (a) Ramanathan, V. Science 1975,190,50-52. (b) Lacis, A.; Hansen, J.; Lee, P.; Mitchell, T.; Lebcdeff, S.Geophys. Res. Lrrr. 1981,8, 1035-38. (4) Montreal Protocol on Substances that Deplete the Ozone Layer, United Nations Environmental Programme, 1987. (5) WMO, Atmospheric Ozone, 1989, Scientific Assessment of Stratospheric Ozone, World Meteorological Organization, Global Ozone Rtsearch and Monitoring Project Report No. 20, Volume 11, Appendix;AFEAS Report, Geneva, Switzerland. (6) Gierczak, T.; Talukdar, R.; Vaghjiani, G. L.; Lovejoy, E. R.; Ravishankara, A. R. J . Geophys. Res. 1991, 96, 5001-5011. (7) Orlando, J. J.; Burkholder. J. B.; McKan, S.A.; Ravishankara, A. R. J . Geophys. Res. 1991,96, 5013-5023. (8) DeMore, W. B.; S. P. Sander, S. P.; Golden, D. M.; Molina, M. J.; Hampson, R. F.; Kurylo, M. J.; Howard, C. J.; Ravishankara, A. R. NASA-JPL, "Chemical kinetics and photochemical data for use in modeling", Evaluation No. 9, NASA Panel for Data Evaluation, JPL Publication 90-1, Jet Laboratory, Pasadena, CA, 1990. (9) Liu, R.; Huie, R. E.; Kurylo, M. J. J. Phys. Chem. 1990, 94, 3247-49. (IO) Brown, A. C.; Canosa-Mas, C. E.; Parr, A. D.; Wayne, R. P. A m " Enoiron. 1990, 24A, 2499.
Q 1991 American Chemical Society
Talukdar et al.
5816 The Journal of Physical Chemistry, Vol. 95, No. 15, 1991
dependence of the cross sections have not been reported. We obtained ultraviolet absorption spectra for this molecule between 190 and 230 nm over the temperature range 203-295 K. The HFCs have no measurable absorption" at wavelengths longer than 180 nm, in agreement with our preliminary measurements. Therefore, photolysis of HFCs in the troposphere or the stratosphere is not an important process and their absorption cross sections were not measured here.
TABLE I: Levels of Impurities in HFC .ad HCFC Samples Used h This study 1. CH2FI (HFC-32) 2. CHiCF, (HFC-143a) 3. CF3CFzH (HFC-125) 4. CHXFCI, (HCFC- 1i l b)
Experimental Section
OH Reaction Studies. The apparatus and the procedures employed in the measurements of kl-k4 are the same as those used by Gierczak et aL6 Detailed descriptions of the experimental methods are given in previous papers (Stimpfle et a1.I2 and Wang et al." on the flow tube-laser magnetic resonance (LMR) method and Vaghjiani and R a ~ i s h a n k a r a on ' ~ pulsed photolysis-laserinduced fluorescence (LIF) method). Therefore, we will present only the aspects necessary to understand the measurements reported here. Discharge Flow-Laser Magnetic Resonance Apparatus. The discharge flow reactor consisted of a 110 cm long, 2.22 cm i.d. Pyrex tube with a 0.95 cm 0.d. movable injector. To reduce wall loss of OH, the inside of the reactor and the outside of the injector were covered with Teflon sleeves. The temperature of the reactor was controlled by flowing through its outer jacket a temperature-regulated fluid from a recirculating reservoir. The OH radicals were produced in a side arm by the fast reaction of H, from a microwave discharge of a H2/He mixture (1%). with an excess of NO2 (2% NO2 in UHP He). The HFCs and HCFCs were added through the movable injector and were always in large excess over OH. The relative concentrations of OH at the end of the flow tube were measured by using laser magnetic resonance spectroscopy. The concentrations of the stable reagents in the flow tube were determined from flow rates (measured with calibrated mass flow meters) and pressure (measured by a capacitance manometer). The linear gas flow velocities ranged from 370 to 2370 cm s-l in 1-3 Torr of helium. The pseudo-first-order rate constants for the decay of O H were obtained by monitoring its LMR signal as a function of reaction time (Le., injector position) at various concentrations of HFCs and HCFCs, the exreagents. The bimolecular rate coefficients were calculated from the slope of pseudo-first-order rate constants plotted against concentrations of the reagents. Pulsed Photolysis-Pulsed Laser-Induced Fluorescence Apparatus. A Pyrex reactor (internal volume of 150 cm3) located in a vacuum housing was maintained at a desired constant temperature by circulating through its outer jacket a fluid from a thermostated bath. The concentrations of all gases were determined by measurements of the mass flow rates and the total pressure. The concentration of CH3CFC12(HCFC-141 b) was also determined from UV absorption at 213.9 nm (Zn lamp). The UV absorption cross section at 213.9 nm used here, (0.945 f 0.080) X 1 O-*O cm2, was determined by slowly flowing the compound in a 100 cm long absorption cell at a known pressure and measuring the absorbance. The above cross section is the average of the values obtained from the two different samples of CH3CFC12 used in this study. This value agrees very well with that measured in our laborator by use of a diode array spectrometer, (0.955 f 0.095) X 10-2gcm2(see below), as well as 0.984 X cm2, measured by Gillotay et al." at 295 K. OH radicals were produced by pulsed photolysis of H 2 0 , between 165 and 185 nm, by a xenon flash lamp, or HONO at 355 nm by the third harmonic of Nd:YAG laser. HONO was prepared in situ in a bubbler by dropping a steady amount of aqueous ( I 1) Gillotay, D.;Simon, P.C.; Brasseur, G. Planef.Space Scl. 1989,37, 105- 108. (12) Stimpfle, R. M.;Perry, R. A,; Howard, C. J. J . Chem. Phys. 1979, 71, 5183-90. (13) Wan& N.S.; Lovejoy, E.R.; Howard, C. J. J. Phys. Chem. I987,91, 5743-49. (14) Vaghjiani, G. L.; Ravishankara, A. R. J . Phys. Chem. 1989, 93,
1948-59.
CHnCFnI 20 ppmv CH;F2 3 20 ppmv CHF, I 30 ppmv; CF3CHzCII 10 ppmv sample 1: CF,CH2CFICHI I 2 ppmv; CiH,FCI 5-4 p&nv~CCizFCF;CII 14 ppmv; CH3CFzCII88 ppmv; CHzCCIF I 1 ppmv; CH2CCIZS 7 ppmv sample 2: CHzCFCI I 1 ppmv; CH3CFzCII1 ppmv; CH,CHFCI d 1 ppmv; C4H5F5I 4 ppmv; C4HzF3CII 1 ppmv; CH2CC12I 1 ppmv; t-CHCICHC1 I 1 ppmv; CH3CC13I 1 PPmv
0.1 M N a N 0 2 solution on to a solution of 10% H2SO4 and the ensuing HONO flowed directly into the reactor.15J6 All experiments were carried out under pseudo-first-order conditions with [HFC or HCFC] >> [OH],, the initial concentration of O H (see tables of results). The temporal profiles of [OH] were determined by measuring the LIF signal at various delay times between the photolysis pulse and the probe laser pulse. These profiles yield k', the first-order decay rate constants. The second-order rate coefficients were obtained from plots of k'vs the concentration of HFC or HCFC. In experiments where HONO was used as the photolytic precursor, contributions of the O H + HONO and O H NO2 reactions (NO2 is an impurity in HONO) to the observed first-order O H decay rate coefficient were determined just before and after measurements in the presence of the reagent. All experiments were camed out under slow flow conditions (linear flow velocities less than 10 cm s-l) with UHP He (>99.999%) as the carrier gas. The total pressure was approximately 100 Torr. UV Absorption Cross-section Measurements. The apparatus used to measure the absorption cross sections was the same as that used in a previous study of HCFCs and Halons in our l a b oratory.'J6 The experimental procedure, data acquisition, and data handling are also described in those papers. Therefore, this information will not be repeated here.
+
Results and Discussion
OH Rate COefiicieaQ. The reactions of HCFCs and HFCs with OH are relatively slow, Le., cm3 molecule-' s-l. Therefore, it is easy to overestimate the rate coefficients if impurities, which react with O H much faster than the reactant, are present in the HFC or HCFC samples. To minimize this error, high-purity samples of the HFCs and HCFCs were used in this kinetic study. Samples of HFC-32,99.99%; HFC-125, >99.99%; and HFC143a, 99.99% were provided by Allied-Signal Corp. Two samples of HCFC- 141b were used. Sample 1, >99.98% pure, was provided by du Pont de Nemours and Co. while sample 2, >99.99%, was from Allied-Signal Corp. Table I lists all the known impurities in the samples as determined by gas chromatography. Another common source of systematic error in measuring small rate coeficients is the reaction of the radical product(s) with the reactant OH. If such a reaction takes place to a significant extent, the rate coefficient for the radical-molecule reaction under consideration would be overestimated. The low initial concentrations of OH (usually
