# Atmospheric chemistry of : Kinetics of the gas-phase reactions with Cl atoms, OH radicals, and O 3 

## M.P. Sulbaek Andersen a , ∗ 

## , E.J.K. Nilsson b 

## , O.J. Nielsen b 

## , M.S. Johnson b 

## , M.D. Hurley c 

## , T.J. Wallington c

a Department of Chemistry, 572 Rowland Hall, University of California, Irvine, CA 92697-2025, USA b Copenhagen Center for Atmospheric Research, Department of Chemistry, University of Copenhagen, Universitetsparken 5, DK-2100 Copenhagen, Denmark c Systems Analytics and Environmental Sciences Department, Ford Motor Company, Mail Drop RIC-2122, Dearborn, MI 48121-2053, USA



### article info Article history:

Keywords: CFC replacement Atmospheric lifetime 

### abstract Long path length Fourier transform infrared (FTIR)–smog chamber techniques were used to study the kinetics of the gas-phase reactions of Cl atoms, OH radicals and O 3 with , in 700 Torr total pressure at 295 ± 2 K. Values of k(Cl + ) = (5.22 ± 0.72) 10

− 11 cm3 molecule

− 1 s − 1 , k(OH + = (4.40 ± 0.38) × 10

− 13 cm3 molecule

− 1 s − 1 and k(O 3 + = (1.46 ± 0.12) × 10

− 21 cm3 molecule

− 1 s − 1 , were established (quoted uncertainties are 2 ; see Experimental section). The IR spectrum of is reported. The atmospheric lif

etime of is determined by reaction with OH radicals and is approximately 26 days. The global warming potential of is approximately 7 for a 100-year time horizon. All rights reserved.

1. Introduction Recognition of the adverse environmental consequences of the release of chlorofluorocarbons (CFCs) into the atmosphere [1,2] hasled to an international effort to replace these compounds with environmentally acceptable alternatives. While safeguarding the ozone layer has been at the center of these efforts, concerns related to global climate change are becoming an increasingly important consideration in the choice of alternative compounds because of their significant atmospheric lifetime and radiative properties. Saturated hydrofluorocarbons (HFCs), such as CH 2 FCF3 (HFC134a), have become widely used CFC replacements. HFCs do notcontain chlorine and therefore do not contribute to chlorinebased catalytic destruction of stratospheric ozone [3] . Unsaturated hydrochlorofluorocarbons are a class of compounds which has been developed as potential replacements for CFCs and saturated HFCs in air conditioning units. Prior to any large-scale industrial use anassessment of the atmospheric chemistry, and hence, environmental impact, of these compounds is needed. To address this need, the atmospheric chemistry of was investigated. The kinetics of the gas-phase reaction of Cl atoms and ozone with were determined at Ford Motor Company (Ford) and at University of Copenhagen (Copenhagen) using smog cham-

∗ Corresponding author. Tel.: +1 562 556 9880; fax: +1 949 824 2905.

E-mail address: mads@sulbaek.dk (M.P.S. Andersen). ber/Fourier transform infrared (FTIR) techniques. The IR spectrum of the title compound and the kinetics of its gas-phase reaction with OH radicals were measured at Ford. Results are reported herein. 2. Experimental Detailed descriptions of the experimental procedures can be found elsewhere [4] . Therefore, only brief descriptions are given in the subsections below. Chlorine atoms were produced by photolysis of molecular chlorine: Cl 2 + hv → Cl + Cl (1) OH radicals were produced by photolysis of CH 3 ONO in the presence of NO in air: CH 3 ONO + hv → CH 3 O + NO (2) CH 3 O + O 2 → HO 2 + HCHO (3) HO 2 + NO → OH + NO 2 (4) In relative rate experiments the following reactions take place: Cl + Reactant → products (5) Cl + Reference → products (6) OH + Reactant → products (7) OH + Reference → products (8)93 It can be shown that
Ln 

### [Reactant] t0 [Reactant] t 

### 

= kReactant kReference Ln 

### [Reference] t0 [Reference] t 

###  (I) where [Reactant] t0 , [Reactant] t , [Reference] t0 , and [Reference] t are the concentrations of reactant and reference at times t0 and t, and kReactant and kReference are the rate constants for the reactant and the reference. Plots of Ln([Reactant] t0 /[Reactant] t ) versus Ln([Reference] t0 /[Reference] t ) should be linear, pass through the origin, and have a slope of kReactant / kReference . The kinetics of the O 3 reaction were studied using an absolute rate method, in which the pseudo first order loss of Reactant was measured in the presence of excess O 3 . A linear plot of the pseudo-first-order rate constants versus the initial O 3 concentration gives a slope of k9 : O 3 + Reactant → products (9) In smog chamber experiments it is important to check for unwanted loss of reactants and products via photolysis, dark chemistry and heterogeneous reactions. Control experiments were performed at Ford and Copenhagen, in which mixtures of reactants (except Cl 2 or CH 3 ONO) in air were subjected to UV irradiation for15 min and reactant/product mixtures obtained after the UV irradiation were allowed to stand in the dark in the chamber for 30 min. There was no observable loss of reactants or reference compoundssuggesting that photolysis, dark chemistry and heterogeneous reactions are not a significant complication in the present work. Unless otherwise stated, quoted uncertainties are two standard deviations from the least squares regressions (forced through zero), together with our estimated uncertainties ( ± 1% of the initial concentration) associated with the IR spectral analysis of the relative reactant and reference concentrations. 2.1. FTIR–smog chamber system at Ford Experiments were performed in a 140-l Pyrex reactor interfacedto a Mattson Sirus 100 FTIR spectrometer. The reactor was surrounded by 22 fluorescent blacklamps (GE F15T8-BL) which were used to photochemically initiate the experiments. CH 3 ONO was synthesized by the drop wise addition of concentrated sulfuric acid to a saturated solution of NaNO 2 in methanol. O 3 was produced from O 2 via silent electrical discharge using a commercial O 3 ozonizer. Samples of were supplied by at purities >99.9%. All other reagents were obtained from commercial sources at purities >99%. Relative rate and absolute rate experimental methods were used as described above. Experiments were conducted in 700 Torr total pressure of air/N 2 diluent at 295 ± 2 K. The loss of and reference compounds was monitored by FTIR spectroscopy using a calibrated analytical path length of 27.1 m and a spectral resolution of 0.25 cm

− 1 . A liquid nitrogen cooled midband MCT detector was used in this study. Infrared spectra were derived from 32 coadded interferograms. Reactant and reference compounds were monitored using absorption features over the following wavenumber ranges (cm

− 1 ): C2 H 2 , 650–800; C2 H4 , 900–1000; O 3 , 980–1070 669, 847, 935, 1311, and 1648.
2.2. FTIR–smog chamber system at University of Copenhagen The Copenhagen experiments were performed in a 100-l quartz reactor equipped with a Bruker IFS 66V FTIR spectrometer. The experimental setup was equipped with a temperature control system that kept the temperature steady at 293 ± 0.5 K. An InSbdetector was chosen for this study. Chlorine atoms were produced by photolysis of Cl 2 according to reaction (1) , using UVA lamps (Osram Eversun L100/79) with the main emissions peak at 368 nm. Ozone was produced using a commercially available ozone generator from O 3 -Technology. The ozone was preconcentrated using a silica gel trap, reducing the amount of O 2 introduced into the chamber. Samples of were supplied by at purities >99.9%. All other reagents were obtained from commercial sources at purities >99%. Relative rate and absolute rate experimental methods were used as described above. The Cl atom relative rate experiments were conducted in 700 Torr total pressure of N 2 diluent. Ozone kinetics experiments were conducted in 700 Torr total pressure of air diluent. The White-type optical system propagating the IR beam through the reaction chamber had a path length of 86 m. Spectra were obtained at a resolution of 0.115 cm

− 1 and were derived from 64 coadded interferograms. Reactant and reference compounds were monitored using absorption over the following wavenumber ranges (cm

− 1 ): C2 H 2 , 3175–3375; C2 H4 , 2940–3040; O 3 , 2950–3065; 3060–3130. FTIR spectra were analyze d using a linear least squares fitting procedure developed by Griffith [5] . Reference spectra of O 3 and HCl were taken from the HITRAN database [6] , and used to improve the spectral fitting forthe larger molecules. For the other components, reference spectra were recorded using the same conditions (700 Torr, 293 ± 0.5 K) employed in the kinetics experiments. 3. Results and discussion 3.1. Relative rate study of k(Cl + ) at Ford and Copenhagen The kinetics of reaction (10) were measured relative to those of reactions (11) and (12) : Cl + → products ( 10) Cl + C2 H4 → products (11) Cl + C2 H 2 → products (12) Initial reaction mixtures employed at Ford consisted of 2.80–3.82 mTorr of 100–102 mTorr Cl 2 , and either 10.0–19.8 mTorr C2 H 4 , or 7.20–10.6 mTorr C2 H 2 , in 700 Torr ofair diluent. Reaction mixtures employed at Copenhagen consisted of 5.6–47.8 mTorr of 100–200 mTorr Cl 2 , and either 4.1–7.1 mTorr of C2 H4 or 3.1–5.0 mTorr of C2 H 2 , in 700 Torr of N 2 diluent. The observed loss of versus those of the reference compounds is plotted in Fig. 1 . Linear least squares analysis of the data in Fig. 1 gives the results shown in Table 1 . The measured rate constant ratios were placed on an absolute scale using k 11 = 9.29 × 10

− 11 cm3 molecule

− 1 s − 1 [7] and k 12 = 5.07 × 10

− 11 cm3 molecule

− 1 s − 1 [7] . The values for k 10 obtained using different experimental setups at Ford MotorCompany and University of Copenhagen were, within the experimental uncertainties, indistinguishable. We derive a final value for k 10 , which is the average of the determinations with error limits which encompass the extremes of the determinations, of k 10 = (5.22 ± 0.72) × 10

− 11 cm3 molecule

− 1 s − 1

. While there have been no previous studies of k9 , we can compare our result with data for the reactions of chlorine atoms with propene and other halogen substituted propenes, presented in Table 2 [8–13] . The reaction of chlorine atoms with propene proceeds primarily via electrophilic addition to the  C C  double bond. Substitution of the electron donating CH 3 group by the electron withdrawing CF3 group would be expected to lower the reactivity of chlorine atoms towards the substituted propene. While the reactivities of t-CF3 CH CHF, E-CF3 CF CHF and Z

CF3 CF CHF, are all comparable and/or within the error limits of the94

Fig. 1. Kinetic data for reactions of Cl atoms with measured at 295 ± 2 K in 700 Torr of air (Ford, A) and in 700 Torr of N 2 (Copenhagen, B). The error bars reflect uncertainty ( ± 1% of the initial concentration) in the determination of the reactant concentrations. individual determinations, inspection of Table 2 reveals a general trend of decreased reactivity with increasing number and electron withdrawing strength of the halogen substituents. 3.2. Relative rate study of k(OH + ) at Ford The kinetics of reaction (13) were measured relative to reactions (14) and (15) : OH + → products (13) OH + C2 H 4 → products (14) Fig. 2. Decay of vs. C2 H 2 (circles) and C2 H4 (squares) in the presence of OH radicals in 700 Torr of air at 295 ± 2 K. The error bars reflect uncertainty ( ± 1% of the initial concentration) in the determination of the reactant concentrations. OH + C2 H 2 → products (15) Initial reaction mixtures consisted of 27.9–32.8 mTorr of 100–330 mTorr CH 3 ONO, and 2.35–7.65 mTorr C2 H4 or 1.47–2.74 mTorr C2 H 2 , and 0–25.4 mTorr NO. Fig. 2 shows the loss of plotted versus loss of the reference compounds. Linear least squares analysis gives k 13 / k 14 = 0.053 ± 0.003 and k 13 / k 15 = 0.506 ± 0.031. Using k 14 = 8.52 × 10

− 12 [14] (atmospheric pressure, 298 K) and k 15 = 8.45 × 10

− 13 [15] gives k 13 = (4.52 ± 0.26) × 10

− 13 and (4.28 ± 0.26) × 10

− 13 cm3 molecule

− 1 s − 1 . Indistinguishable values of k 13 were obtained using the two different references. The fact that consistent values of k 13 were derived from experiments using different reference compounds suggests the absence of significant systematic errors in the present work. We choose to cite a final value which is the average of the individual determinations with error limits which encompass the extremes of the determinations, yielding k 13 = (4.40 ± 0.38) × 10

− 13 cm3 molecule

− 1 s − 1 . No result has been reported for k 13 previously. Kinetic data for the reactions of hydroxyl radicals with propene and fluorinated propenes are presented in Table 2 [8–13,16–20]

.Based on the structural similarity, one might expect the reactivity of towards OH radicals to be comparable to that of t-CF3 CH CHF. Interestingly, the reactivity of is substantially lower (roughly a factor of 2). As with chlorine atoms, the reaction with hydroxyl radicals proceeds mainly via electrophilic addition to the  C C  double bond [14] . Hence, the reactivity towards OH radicals is expected to decrease with the introduction of electron withdrawing substituents such as CF3 – groups, chlorine and fluorine. Inspection of Table 2 reveals that thereactivity of OH radicals towards propene is greater than the fluorinated propenes and that there is a general trend toward lower

Table 1 Results of the Cl atom relative rate experiments Setup Reference k /kReference kReference ( × 10
− 11 cm3 molecule

− 1 s − 1 ) kCl + ( × 10
− 11 cm3 molecule

− 1 s − 1 ) Ford (air) C2 H4 0.55 ± 0.04 9.29 [7] (5.11 ± 0.37) Ford (air) C2 H 2 1.03 ± 0.06 5.07 [7] (5.22 ± 0.30) Copenhagen (N 2 ) C2 H4 0.59 ± 0.05 9.29 [7] (5.48 ± 0.46) Copenhagen (N 2 ) C2 H 2 1.00 ± 0.09 5.07 [7] (5.07 ± 0.46)95 Table 2 Rate constants (cm3 molecule

− 1 s − 1 ) for reactions of Cl, OH, and O 3 with and analogous halogenated propenes measured at 295 ± 3 K (where multiple determinations exist, the value given is the average of the determinations with error limits that encompass the extremes of the individual determinations) Compound kCl kOH kO3 CF3 CF CF2 (2.7 ± 0.3) × 10
− 11 [11] (2.3 ± 0.7) × 10
− 12 [11,16–19] (6.2 ± 1.5) × 10
− 22 [21] Z- CF3 CF CHF (4.36 ± 0.48) × 10
− 11 [12] (1.26 ± 0.20) × 10
− 12 [12,20] (1.45 ± 0.15) × 10
− 21 [12] E- CF3 CF CHF (5.00 ± 0.56) × 10
− 11 [12] (2.15 ± 0.23) × 10
− 12 [12] (1.98 ± 0.15) × 10
− 20 [12] t- CF3 CH CHF (4.64 ± 0.59) × 10
− 11 [13] (9.25 ± 1.72) × 10
− 13 [13] (2.81 ± 0.21) × 10
− 21 [13] (5.22 ± 0.72) × 10
− 11 a (4.40 ± 0.38) × 10
− 13 a (1.46 ± 0.12) × 10
− 21 a CF3 CF CH 2 (7.03 ± 0.59) × 10
− 11 [10] (1.10 ± 0.22) × 10
− 12 [10,17,20] (2.77 ± 0.21) × 10
− 20 [10] CF3 CH CH 2 (9.07 ± 1.08) × 10
− 11 [8] (1.45 ± 0.33) × 10
− 12 [8,17] (3.5 ± 0.3) × 10
− 19 [8] CH 3 CH CH 2 (2.7 ± 0.7) × 10
− 10 [9] (2.9 ± 0.3) × 10
− 11 [9] (1.0 ± 0.1) × 10
− 17 [9]

a Determined in this work.
reactivity with increasing fluorination. However, in contrast to this general trend, perfluoropropene is significantly more reactive than the less fluorinated propenes. Furthermore, in contrast to a simpletrend of decreasing reactivity with increasingly electron withdrawing substituents, as noted above, CF3 CH CHF is more reactive than It has recently been suggested that hydrogen bonding b etween OH and fluorine could lead to the observed enhancement of the reactivity of OH radicals towards the highly fluorinated propenes [13] . Computational work would be useful to investigate this suggestion. 3.3. Absolute rate study of k(O3 + at Ford and Copenhagen The kinetics of reaction (16) were studied by observing the decay of l when exposed to O 3 in excess in the reaction chambers at Ford and Copenhagen: O 3 + → products ( 16)An OH radical scavenger (cyclohexane) was added to the reaction mixtures to avoid potential problems associated with the loss of via reaction with OH radicals formed in reaction (1 6) [14] . Initial concentrations of reactants at Copenhagen were 6.9–14.2 mTorr 15.6–20.3 mTorr cyclohexane, and 3.42–4.36 Torr O 3 in 700 Torr of air diluent. Initial concentrations of reactants at Ford were 5.9–6.3 mTorr 1 4.0–30.6 mTorr cyclohexane, and 414–2070 mTorr O 3 in 700 Torr of air diluent. Variation of the [cyclohexane]/ ] ratio over the range 2.2–5.2 had no discernable effect on the observed decay of suggesting that loss via reaction with OH radicals is not a significant complication. The loss of followed pseudo first-order kinetics in all experiments (see insert in Fig. 3 ). A plot of the pseudo first-order decay of versus O 3 concentration is shown in Fig. 3 . A linear least squares fit to the data gives k 16 = (1.46 ± 0.10) × 10

− 21 cm3 molecule

− 1 s − 1

. We choose to cite a final value for k 16 , with error limits whichinclude two standard deviations from the least squares regression and a propagated 5% uncertainty in the O 3 calibration, of k 16 = (1.46 ± 0.12) × 10

− 21 cm3 molecule

− 1 s − 1

. The result is compared to the reactivity of O 3 towards propene and analogous fluorine substituted propenes in Table 2 [8–10,12,13,21] . In its reaction with O 3 , t-CF3 CH CHF is less reactive than its non-substituted counterpart by a factor of 6850. The reaction occurs by electrophilic addition to the  C C  double bond. When compared with Cl atoms and OH radicals, O 3 has the lowest reactivity and hence, is most sensitive to the presence of electron withdrawing constituents. The existing literature shows a broad trend of decreasing reactivity towards O 3 with increasing number and electron withdrawing strength of the halogen constituents. However, there are exceptions to this broad trend as evident from inspection of the data in Table 2 . The present work shows that has a remarkably low reactivity towards O3 ; lower than the corresponding fluorinated compound t-CF3 CH CHF, lower than E-CF3 CF CHF, and comparable to Z-CF3 CF CHF. It is clear thatdecreasing reactivity with increasingly electron withdrawing substituents is not entirely sufficient to explain the observed reactivity trend. Computational work would be of interest to investigate thefactors involved in determining the trend of reactivity of the halopropenes towards ozone. 3.4. Infrared spectrum of The IR spectrum of measured in the present work in 700 Torr of air diluent at 295 ± 2 K is shown in Fig. 4

. has IR features at 669, 847, 93 5, 1152, 1311 and 1648 cm

− 1 . The total (600–1800 cm

− 1 ) integrated IR absorption cross-section of is (1 .74 ± 0.20) × 10

− 16 cm molecule

− 1 . Individual IR band absorption cross-sections are 6.22 × 10

− 18 (600–690 cm

− 1 ), 2.15 × 10

− 17 (750– 1030 cm

− 1 ), 1.27 × 10

− 16 (1030–1350 cm

− 1 ) and 1.78 × 10

− 17 (1600–1790) cm molecule

− 1 . There are no literature IR data for to compare with our result.

Fig. 3. Pseudo-first order loss o vs. O 3 concentration. Data obtainedat Ford Motor Company and at University of Copenhagen are shown with open symbols and gray symbols, respectively. The insert shows the decay of when exposed to 414 mTorr (open circles), 911 mTorr (open upward triangles), 1.47 Torr (open downward triangles), 2.07 Torr (open squares), 3.42 Torr (gray circles) and 4.36 Torr (gray triangles) of O 3 in 700 Torr of air at 295 ± 2 K.96

Fig. 4. IR spectrum of in 700 Torr of air, 295 ± 2 K (base e). 4. Implications for atmospheric chemistry The present work improves our understanding of the atmospheric chemistry of Cl atoms and OH radicals react with with rate coefficients of (5.22 ± 0.72) × 10

− 11 and (4.40 0.38) 10

− 13 cm3 molecule

− 1 s − 1 , respectively. O 3 reacts with with a rate constant of (1 .46 ± 0.12) × 10

− 21 cm3 molecule

− 1 s − 1

. is not expected to undergo wet or dry deposition or photolysis [17] to an appreciable extend. Cl atoms are not present in the atmosphere in sufficient quantities to have a significant influence on the lifetime of The value of k(OH + derived in the present work can b e used to provide an estimate of the atmospheric lifetime of Using a global weighted-average OH concentration of 1 .0 10 6 molecules cm

− 3 [22] leads to an estimated lifetime of with respect to reaction with OH radicals of 26 days. In a similar fashion our value for k(O 3 + derived above can b e combined with the global background concentration of O 3 of approximately 35 ppb [23] to provide an estimate of the atmospheric lifetime of with respect to reaction with O 3 of 25 years. Hence we conclude that the atmospheric lifetime of is determined by reaction with OH radicals and is approximately 26 days. The approximate nature of this lifetime estimate should be stressed; the average daily concentration of OH radicals in the atmosphere varies significantly with both location and season [24] . The values above are estimates of global average lifetimes; local lifetimes could be different from those quoted above. Using the method outlined by Pinnock et al. [25] , the IR spectra of shown in Fig. 4 and the IR spectrum of CFC-11 [26] we calculate instantaneous forcings for and CFC-11 of 0.2 14 and 0.26W m

− 2 ppb

− 1 , respectively. Values of the halocarbon global warming potential, HGWP [27] , for (relative to CFC-11) can then be estimated using the expression (II) : HGWP = 

### IF IFCFC-11 

###   MCFC-11  CFC-11 M 

### 

× 

### 1

− exp(

− t/ ) 1

− exp(

− t/ CFC-11 ) 

###  (II)
where IF , IFCFC-11 , M , MCFC-11 ,  and  CFC-11 are the instantaneous forcings, molecular weights and atmospheric lifetimes of and CFC-11, and t is the time horizon over which the forcing is integrated. Using  = 26 days and CFC-11 = 45 years [28] we estimate that the HGWP of relative to CFC-11 is 3.8 × 10

− 3 for a 20-year horizon and 1.5 × 10

− 3 for a 100-year time horizon, respectively. Relative to CO 2 , the GWP of CFC-11 on 20- and 100-year time horizons are 6730 and 4750, respectively [28] . Therefore, relative to CO 2 , the GWP of is approximately 26 for a 20- year horizon and 7 for a 100-year time horizon. has a negligible global warming potential and will not make a significant contribution to radiative forcing of climate change. It has been estimated that compounds with a tropospheric lifetime of 25 days, uniform landmass emissions, approximately the same mass as CFC-11 and containing one chlorine atom have an ozone depletion potential of approximately 0.003 [29] . The contribution of to stratospheric ozone depletion is expected to be negligible given its very short-atmospheric lifetime. Acknowledgements We thank ( for providing the high-purity samples of used in this study. OJN, MSJ and EJKN acknowledge financial support from the Danish Natural Science Research Council and the Villum Kann Rasmussen Foundation for the Copenhagen Center for Atmospheric Research (CCAR). MPSA acknowledges financial support from the Comer Foundation and thanks F.S. Rowland and D.R. Blake (University of California, Irvine) for helpful discussions. Appendix A. Supplementary data Supplementary data associated with this article can be found, in the online version, at
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