domain: chim phys

text:
Yuri Bedjanian
email: yuri.bedjanian@cnrs-orleans.fr

Temperature-Dependent Kinetic Study of the Reaction of Hydroxyl Radical with Hydroxyacetone

The kinetics of the reaction of OH radicals with hydroxyacetone has been investigated as a function of temperature at (2.0 -2.1) Torr total pressure of helium and over an extended temperature range, T = 250 -830 K, and as a function of pressure at T = 301 K in the pressure range 1.0-10.4 Torr. The rate constant of the reaction OH + CH 3 C(O)CH 2 OH  products (1) was measured using both absolute (from the kinetics of OH consumption in excess of hydroxyacetone) and relative rate method, k 1 = 4.7×10 -22 ×T 3.25 exp(1410/T) cm 3 molecule -1 s -1 at T = 250-830 K. The present data combined with selected previous temperature dependent studies of reaction (1) yield k 1 = 4.4×10 -20 ×T 2.63 exp(1110/T) cm 3 molecule -1 s -1 , which is recommended from the present work at T = 230-830K (with conservative uncertainty of 20% at all temperatures). k 1 was found to be pressure independent in the pressure range from 1.0 to 10.4 Torr of He at T = 301 K. The present results are compared with previous experimental and theoretical data.

INTRODUCTION

Hydroxyacetone (CH 3 C(O)CH 2 OH) is the simplest representative of hydroxyketones, chemical species belonging to a large family of Oxygenated Volatile Organic Compounds (OVOC) playing an important role in the tropospheric chemistry. [START_REF] Calvert | Mechanisms of Atmospheric Oxidation of the Oxygenates[END_REF] Hydroxyacetone (HAC) is one of the most abundant OVOC in the atmosphere due to important biogenic emissions [START_REF] Christian | Comprehensive Laboratory Measurements of Biomass-Burning Emissions: 1. Emissions from Indonesian, African, and other Fuels[END_REF] and in situ production in atmospheric oxidation of isoprene [START_REF] Karl | Rapid Formation of Isoprene Photo-Oxidation Products Observed in Amazonia[END_REF][START_REF] Paulot | Isoprene Photooxidation: New Insights Into the Production of Acids and Organic Nitrates[END_REF][START_REF] Galloway | Yields of Oxidized Volatile Organic Compounds During the OH Radical Initiated Oxidation of Isoprene, Methyl Vinyl Ketone, and Methacrolein under High NO X Conditions[END_REF] and its major OH-initiated oxidation product, methacrolein. [START_REF] Galloway | Yields of Oxidized Volatile Organic Compounds During the OH Radical Initiated Oxidation of Isoprene, Methyl Vinyl Ketone, and Methacrolein under High NO X Conditions[END_REF][START_REF] Crounse | Atmospheric Fate of Methacrolein. 1. Peroxy Radical Isomerization Following Addition of OH and O 2[END_REF] The atmospheric degradation of hydroxyketones is triggered by their photolysis and reactions with atmospheric oxidants (OH, NO 3 ,O 3 , Cl), mainly with OH radical. [START_REF] Calvert | Mechanisms of Atmospheric Oxidation of the Oxygenates[END_REF] The understanding of the atmospheric chemistry of hydroxyketones requires information on their major photo-oxidation routs, including kinetics of their reactions with the atmospheric oxidants. Reactions of OH radicals with hydroxyketones proceed through a complex kinetic mechanism leading to a non-Arrhenius behavior of the reaction rate constant, [START_REF] Bedjanian | Temperature-Dependent Rate Constant for the Reaction of Hydroxyl Radical with 3-Hydroxy-3-methyl-2-butanone[END_REF][START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF] and in this respect, are of great interest for the theory. The existing data base on the reactions of hydroxyketones with OH is mainly limited to a number of relative rate measurements at room temperature. [START_REF] Calvert | Mechanisms of Atmospheric Oxidation of the Oxygenates[END_REF] In our recent paper, we have reported the temperature dependence of the rate constant for the reaction of OH with 3-hydroxy-3-methyl-2-butanone in the temperature range 278-830K. [START_REF] Bedjanian | Temperature-Dependent Rate Constant for the Reaction of Hydroxyl Radical with 3-Hydroxy-3-methyl-2-butanone[END_REF] The reaction of OH with hydroxyacetone OH + CH 3 C(O)CH 2 OH  products (1) has been quite extensively studied both experimentally [START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF][START_REF] Dagaut | Kinetic Measurements of the Gas-Phase Reactions of Hydroxyl Radicals With Hydroxy Ethers, Hydroxy Ketones, and Keto Ethers[END_REF][START_REF] Orlando | The Rate and Mechanism of the Gas-Phase Oxidation of Hydroxyacetone[END_REF][START_REF] Chowdhury | ArF Laser Photodissociation Dynamics of Hydroxyacetone: LIF Observation of OH and Its Reaction Rate with the Parent[END_REF][START_REF] Butkovskaya | Mechanism of the OH-Initiated Oxidation of Hydroxyacetone over the Temperature Range 236-298 K[END_REF][START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF][START_REF] Baasandorj | Experimental and Theoretical Studies of the Kinetics of the OH + Hydroxyacetone Reaction As a Function of Temperature[END_REF] and theoretically. [START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF][START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF][START_REF] Baasandorj | Experimental and Theoretical Studies of the Kinetics of the OH + Hydroxyacetone Reaction As a Function of Temperature[END_REF][START_REF] Galano | Theoretical Study on the Reaction of Tropospheric Interest: Hydroxyacetone + OH. Mechanism and Kinetics[END_REF] Theoretical calculations [START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF][START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF][START_REF] Baasandorj | Experimental and Theoretical Studies of the Kinetics of the OH + Hydroxyacetone Reaction As a Function of Temperature[END_REF][START_REF] Galano | Theoretical Study on the Reaction of Tropospheric Interest: Hydroxyacetone + OH. Mechanism and Kinetics[END_REF] showed that reaction (1) proceeds through pre-reactive hydrogen-bonded complexes OH + HAC  OHHAC*  OHHAC  products where OHHAC* and OHHAC denote excited and collisionally thermalized OHhydroxyacetone complexes, respectively. The reaction of hydroxyacetone with OH radical is supposed to proceed mainly by H-atom abstraction from CH 2 group, Despite the relatively large number of experimental studies of reaction (1), [START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF][START_REF] Dagaut | Kinetic Measurements of the Gas-Phase Reactions of Hydroxyl Radicals With Hydroxy Ethers, Hydroxy Ketones, and Keto Ethers[END_REF][START_REF] Orlando | The Rate and Mechanism of the Gas-Phase Oxidation of Hydroxyacetone[END_REF][START_REF] Chowdhury | ArF Laser Photodissociation Dynamics of Hydroxyacetone: LIF Observation of OH and Its Reaction Rate with the Parent[END_REF][START_REF] Butkovskaya | Mechanism of the OH-Initiated Oxidation of Hydroxyacetone over the Temperature Range 236-298 K[END_REF][START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF][START_REF] Baasandorj | Experimental and Theoretical Studies of the Kinetics of the OH + Hydroxyacetone Reaction As a Function of Temperature[END_REF] the value of the reaction rate constant at room temperature, temperature dependence of the rate constant and its potential pressure dependence is still an open question. The results of earlier room temperature measurements of k 1 [START_REF] Dagaut | Kinetic Measurements of the Gas-Phase Reactions of Hydroxyl Radicals With Hydroxy Ethers, Hydroxy Ketones, and Keto Ethers[END_REF][START_REF] Orlando | The Rate and Mechanism of the Gas-Phase Oxidation of Hydroxyacetone[END_REF][START_REF] Chowdhury | ArF Laser Photodissociation Dynamics of Hydroxyacetone: LIF Observation of OH and Its Reaction Rate with the Parent[END_REF][START_REF] Butkovskaya | Mechanism of the OH-Initiated Oxidation of Hydroxyacetone over the Temperature Range 236-298 K[END_REF] in the pressure range from 10 to 750 Torr, employing both absolute [START_REF] Dagaut | Kinetic Measurements of the Gas-Phase Reactions of Hydroxyl Radicals With Hydroxy Ethers, Hydroxy Ketones, and Keto Ethers[END_REF][START_REF] Chowdhury | ArF Laser Photodissociation Dynamics of Hydroxyacetone: LIF Observation of OH and Its Reaction Rate with the Parent[END_REF][START_REF] Butkovskaya | Mechanism of the OH-Initiated Oxidation of Hydroxyacetone over the Temperature Range 236-298 K[END_REF] and relative rate [START_REF] Orlando | The Rate and Mechanism of the Gas-Phase Oxidation of Hydroxyacetone[END_REF] methods, are consistent with each other and fall in the range (2.5-3.5)10 -12 cm 3 molecule -1 s -1 with an average value of k 1 (±2) = (3.0±0.8)10 -12 cm 3 molecule -1 s -1 at T = 298 K. In two subsequent temperature dependent studies of reaction

(1), negative and positive dependence of k 1 on temperature were reported by Dillon et al. at 60

Torr [START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF] and Baasandorj et al. at 2-5 Torr of He, [START_REF] Baasandorj | Experimental and Theoretical Studies of the Kinetics of the OH + Hydroxyacetone Reaction As a Function of Temperature[END_REF] respectively. Besides, the value of k 1 at T = 298 K measured by Dillon et al. [START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF] was by a factor of 2 higher compared with other studies.

Finally, in the most recent temperature dependent study of reaction (1), Vu et al. [START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF] have observed a transition in temperature dependence of k 1 from negative at T < 380 K to positive at T = 410-500 K. Moreover, the authors observed a pronounced decrease of k 1 at total pressures lower than ∼45 Torr at T = 301 K. Based on the experimental observations, Vu et al. [START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF] proposed a reaction mechanism, which, according to the authors, reconciles the disparate k 1 high pressure data of Dillon et al. [START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF] and those of Baasandorj et al. [START_REF] Baasandorj | Experimental and Theoretical Studies of the Kinetics of the OH + Hydroxyacetone Reaction As a Function of Temperature[END_REF] measured at low pressures.

The objective of the present low pressure study of reaction (1) was to measure the reaction rate constant in an extended temperature range (250-830 K) and verify the possible pressure dependence of k 1 at low pressures (1-10 Torr of He) in comparison with previous experimental data and theoretical predictions.

EXPERIMENTAL

The experimental setup consisted of a discharge flow reactor combined with a modulated molecular beam mass spectrometer. [START_REF] Bedjanian | Temperature-Dependent Rate Constant for the Reaction of Hydroxyl Radical with 3-Hydroxy-3-methyl-2-butanone[END_REF][START_REF] Morin | Experimental Study of the Reactions of OH Radicals with Propane, n-Pentane, and n-Heptane over a Wide Temperature Range[END_REF][START_REF] Bedjanian | Gas-Phase Reaction of Hydroxyl Radical with p-Cymene over an Extended Temperature Range[END_REF][START_REF] Braure | Experimental Study of the Reactions of Limonene with OH and OD Radicals: Kinetics and Products[END_REF] We have used two flow reactors. The first one, used at low temperatures (250 -336 K), consisted of a Pyrex tube (45 cm length and 2.4 cm i.d.)

with a jacket for the thermostated liquid circulation (ethanol). The walls of this low temperature reactor were coated with halocarbon wax in order to reduce the heterogeneous loss of OH. The second, high temperature, reactor used at T = 310 -830 K, consisted of an uncoated electrically heated Quartz tube (45 cm length and 2.5 cm i.d.) with water-cooled attachments (Figure S1, Supporting Information, SI). [START_REF] Morin | Experimental Study of the Reactions of OH Radicals with Propane, n-Pentane, and n-Heptane over a Wide Temperature Range[END_REF] Hydroxyl radicals were generated inside the movable injector in reaction of H atoms with excess NO 2 ([NO 2 ]  10 13 molecule cm -3 ), hydrogen atoms being produced in a microwave discharge of H 2 /He mixtures (Figure S1, SI):

H + NO 2  OH + NO (2)
k 2 = (1.47 ± 0.26)  10 -10 cm 3 molecule -1 s -1 (T = 195 -2000). [START_REF] Su | Rate Constants, 1100 ≤ T ≤ 2000 K, for H + NO 2 → OH + NO Using Two Shock Tube Techniques: Comparison of Theory to Experiment †[END_REF] OH radicals were scavenged with an excess of Br 2 (added in the end of the reactor 5 cm upstream of the sampling cone, as shown in Figure S1, SI) through reaction 3 and detected at m/z = 96/98 (HOBr + ):

OH + Br 2  HOBr + Br (3) -11 exp(207/T) cm 3 molecule -1 s -1 (T = 220 -950). [START_REF] Bedjanian | Temperature-Dependent Rate Constant for the Reaction of Hydroxyl Radical with 3-Hydroxy-3-methyl-2-butanone[END_REF] The absolute concentration of OH radicals was determined from the consumed fraction of papers from this group. [START_REF] Bedjanian | Kinetic Study of the Reactions of Br 2 with OH and OD[END_REF][START_REF] Bedjanian | Kinetic Study of OH + OH and OD + OD Reactions[END_REF] The hydroxyacetone was introduced into the reactor by flowing the bath gas (He) over liquid hydroxyacetone and was detected by mass spectrometry at its parent peak at m/z = 74 (CH 3 C(O)CH 2 OH + ). The water impurity in the gaseous samples of hydroxyacetone, directly monitored at m/z =18 (H 2 O + ), was always less than 2.5%. The absolute concentration of hydroxyacetone was determined through injection of known amounts (0.2-1.0 μL) of HAC into a helium flow in the reactor and recording its parent peak intensity at m/z = 74. The integrated area of the mass spectrometric signal, corresponding to known total number of HAC molecules injected into the reactor, allowed the determination of the calibration factor.

k 3 = 2.1610
The impact of the possible loss of HAC on the walls of the reactor on the results of the calibration experiments was considered as insignificant for the following reasons. Firstly, in the kinetic measurements, hydroxyacetone was introduced into the reactor through the movable injector (Figure S1, SI) and its kinetics could be directly monitored: we have not observed any measurable loss of HAC on the walls of the reactor in the whole temperature range of the study. Secondly, generally, calibration experiments consisted in several (4-5) successive injections of HAC: the integrated area of the mass spectrometric signal of HAC (normalized by the amount of HAC injected) from different injections were similar within 10%. In addition, the calibration factors determined before the kinetic measurements (that is in the reactor not treated with hydroxyacetone and reaction products) were similar to those measured at the end of the kinetic experiments. Finally, the relationships between the HAC calibration factors and those of the non-sticky species (Br 2 , NO 2 ) determined at high and low temperatures (where the probable wall loss of HAC should be more effective) and in two reactors with different surfaces were similar. Absolute concentrations of NO 2 , Br 2 and n-heptane were calculated from their flow rates obtained from the measurements of the pressure drop of their mixtures in He stored in calibrated volume flasks.

The purities of the gases used were as follows: He >99.9995% (Alphagaz), passed through liquid nitrogen trap; H 2 > 99.998% (Alphagaz); Br 2 >99.99% (Aldrich); NO 2 > 99% (Alphagaz); n-heptane >99% (Aldrich); Hydroxyacetone >95% (Alfa Aesar), degassed before use. 1. S3 (SI). The diffusion corrections, [START_REF] Kaufman | Kinetics of Elementary Radical Reactions in the Gas Phase[END_REF] applied to all the measured values of k 1 ' using diffusion coefficient D 0 (OH-He) = 660 × (T/298) 1.85 Torr cm -2 s -1 , [START_REF] Ivanov | HO 2 , and Ozone Gaseous Diffusion Coefficients[END_REF] were generally within a few

RESULTS AND DISCUSSION

[OH] (relative units)

percent reaching up to 17% in several kinetic runs. One can note that in the experiments at T = 500 K, the variation of the initial concentration of OH from 1.4 to 3.2×10 11 molecule cm -3 had no impact on the measured k 1 '. This indicates that the possible impact of secondary reactions of OH radicals on the measurements of k 1 was insignificant due to their low initial concentrations used. The absolute measurements of k 1 have been carried out in an uncoated quartz reactor, except two temperatures, 298 and 336K, where Pyrex reactor coated with halocarbon wax was used. In the quartz reactor, the intercepts of the linear plots, like those shown in Figure 2, k w , were found to decrease with increasing temperature (from 60 to 10 s -1 between 330 and 830K)

and to be significantly higher (particularly at low temperatures) compared with the rate of OH decay measured in the absence of hydroxyacetone in the reactor (k w 10 s -1 , independent of temperature). Similar behavior was observed in a Pyrex reactor coated with halocarbon wax: the intercepts observed at T = 336 and 298K were found to be about 20 s -1 (Figure S3, SI) and increased at lower temperatures. Baasandorj et al. [START_REF] Baasandorj | Experimental and Theoretical Studies of the Kinetics of the OH + Hydroxyacetone Reaction As a Function of Temperature[END_REF] reported a similar effect in their study of reaction (1) in a reactor coated with halocarbon wax: upon addition of hydroxyacetone some experimental pseudo-first-order decays of OH were nonlinear and large positive intercepts on the second order plots were observed. The observation was tentatively attributed to the heterogeneous reaction between OH radicals and hydroxyacetone adsorbed to the walls of the reactor. In order to avoid the heterogeneous complications, the authors 14 carried out the measurements with addition of oxygen ((2-5)×10 15 cm -3 ) which appeared to minimize the hydroxyacetone-catalyzed heterogeneous loss of OH. In the present study, we have not observed significant influence of oxygen addition on the kinetics of OH, and the reaction rate constant was determined from the slope of the linear fit to the experimental data (straight lines in Figures 2 andS3), that is under assumption that heterogeneous contribution to the loss of OH radicals is independent of hydroxyacetone concentration in the gas phase and considering that the intercepts are relatively low compared with the measured maximum values of k 1 ′. The lowest temperature in the absolute measurements of k 1 was limited to 298 K. It can be noted that very close values of k 1 (see Table 1) were measured at T = 330K in the uncoated quartz reactor (with intercept of  60 s -1 ) and at T = 336K in the halocarbon wax coated reactor (with intercept of  20 s -1 ). The results of the absolute measurements of k 1 are displayed in Table 1.

The combined uncertainty on k 1 was estimated to be approximately 15% by adding in quadrature the uncertainties on the absolute concentration of hydroxyacetone (10%), on the measurements of the flows (3%), pressure (2%), temperature (1%) and statistical error (≤ 5%). 2) were calculated with k 3 = 2.1610 -11 exp(207/T) cm 3 molecule -1 s -1 which is well established in the temperature range 220 -950 K. [START_REF] Bedjanian | Temperature-Dependent Rate Constant for the Reaction of Hydroxyl Radical with 3-Hydroxy-3-methyl-2-butanone[END_REF] One can note that two linear plots in Figure 3, corresponding to the experimental data obtained at T = 750 K with different concentrations of Br 2 , have similar slopes (i.e. provide similar k 1 /k 3 ratios) and different intercepts in agreement with equation (I). In another series of experiments, the rate constant of reaction (1) was measured at two temperatures, T = 290 and 320 K, and P = 2.1 Torr using the traditional relative rate method, which consisted of the measurements of the consumption of hydroxyacetone and reference compound, n-heptane, simultaneously present in the reactor, in reactions with OH:

Rate constant of reaction

OH + HAC  Products (1) OH + n-heptane  Products (4)
The relative consumptions of HAC and n-heptane are defined by the rate constants of their reactions with OH:

(II)

where the expressions under the logarithm are the ratios of the compound concentration in the absence of to that in the presence of OH radicals for a given reaction time.

In the experiments, the initial concentrations of HAC and n-heptane were in the range (4 -6) and (1 -2)× 10 12 molecule cm -3 , respectively. The concentration of OH radicals, formed through reaction ( 2) with [NO 2 ] 5×10 13 molecule cm -3 , was varied being 10 13 molecule cm - 3 at maximum. Hydroxyacetone and n-heptane were monitored by mass spectrometry at their parent peaks at m/z = 74 and 100, respectively. It was verified that there was no interference between the mass spectra of these two species and that the products of their reactions with OH did not contribute to these peaks. Secondary reactions of HAC and n-heptane with organic radicals produced in reactions (1) and ( 4 with uncertainty of 15%. [START_REF] Morin | Experimental Study of the Reactions of OH Radicals with Propane, n-Pentane, and n-Heptane over a Wide Temperature Range[END_REF] With k 4 calculated through this expression and k 1 /k 4 (±1) = 1.093±0.006 and 0.799±0.006 derived from the slopes of the straight lines in Figure 4, k 1 = 6.7110 -12 and 5.5910 -12 cm 3 molecule -1 s -1 at T = 290 and 320 K, respectively, with estimated total uncertainty of 20%, including that on the rate constant of the reference reaction. Similar and independent of pressure values of k 1 were measured in the experiments using reaction of OH + n-heptane as a reference (Figure 6). The values of k 1 /k 4 = 0.965, 0.928 and 0.911 derived from the data in Figure 6 result in k 1 = 6.22, 5.98 and 5.8710 -12 cm 3 molecule - All the data obtained for k 1 at T = 301 K and different pressures are displayed in Figure 7 in line with those reported by Vu et al. [START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF] The present measurements do not show any pressure dependence in the pressure range from 1 to 10.4 Torr and are very close to k 1 data measured by Vu et al. [START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF] at pressures 30-70 Torr of He. Although our data do not confirm a trend of decrease of k 1 with decreasing pressure observed by Vu et al., [START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF] they do not exclude the hypothetical dependence of k 1 on pressure at higher pressures if k 1 is in a low pressure limit in the pressure range (1 -10 Torr) of the present study. Unfortunately, the configuration of the use of the electron impact ionization mass spectrometer in the present work did not allow kinetic measurements at higher pressures.

It can be noted that in another low pressure study of Baasandorj et al., [START_REF] Baasandorj | Experimental and Theoretical Studies of the Kinetics of the OH + Hydroxyacetone Reaction As a Function of Temperature[END_REF] the reaction rate constant was also found to be independent of pressure between 2 and 5 Torr of He and to be in good agreement with room temperature measurements of k 1 by Dagaut et al. [START_REF] Dagaut | Kinetic Measurements of the Gas-Phase Reactions of Hydroxyl Radicals With Hydroxy Ethers, Hydroxy Ketones, and Keto Ethers[END_REF] [START_REF] Butkovskaya | Mechanism of the OH-Initiated Oxidation of Hydroxyacetone over the Temperature Range 236-298 K[END_REF] However, these data cannot be used as an argument against the dependence of the rate constant on pressure, given the expected weak pressure dependence and a relatively large scatter in the existing experimental data (up to a factor of 2 at room temperature, as discussed below). K. The values of k 1 reported by Dillon et al., [START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF] Vu et al. [START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF] and in the present work at T = 298 K are approximately a factor of two higher compared with those reported in earlier room temperature studies [START_REF] Dagaut | Kinetic Measurements of the Gas-Phase Reactions of Hydroxyl Radicals With Hydroxy Ethers, Hydroxy Ketones, and Keto Ethers[END_REF][START_REF] Orlando | The Rate and Mechanism of the Gas-Phase Oxidation of Hydroxyacetone[END_REF][START_REF] Chowdhury | ArF Laser Photodissociation Dynamics of Hydroxyacetone: LIF Observation of OH and Its Reaction Rate with the Parent[END_REF][START_REF] Butkovskaya | Mechanism of the OH-Initiated Oxidation of Hydroxyacetone over the Temperature Range 236-298 K[END_REF] and by Baasandorj et al. [START_REF] Baasandorj | Experimental and Theoretical Studies of the Kinetics of the OH + Hydroxyacetone Reaction As a Function of Temperature[END_REF] Dillon et al. [START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF] and Baasandorj et al. [START_REF] Baasandorj | Experimental and Theoretical Studies of the Kinetics of the OH + Hydroxyacetone Reaction As a Function of Temperature[END_REF] carried out a detailed analysis of the possible errors due to the measurements of the absolute concentrations of hydroxyacetone, including uncertainties on its absorption cross section, presence of impurities and the regeneration of OH; nevertheless, the cause of the disagreement between two sets of k 1 measurements remains unclear. As one can see in Figure 8, the opposite trend in temperature dependence was observed by Dillon et al., [START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF] Vu et al. [START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF] and in the present study compared to Baasandorj et al., [START_REF] Baasandorj | Experimental and Theoretical Studies of the Kinetics of the OH + Hydroxyacetone Reaction As a Function of Temperature[END_REF] who observed a positive temperature dependence of k 1 = (1.88±0.75)×10 -11 exp(-545±60/T) cm 3 molecule -1 s -1 at T = 280-350 K using a discharge-flow system coupled with resonance fluorescence detection of the OH radical. The reason for this discrepancy also remains unclear. Baasandorj et al. [START_REF] Baasandorj | Experimental and Theoretical Studies of the Kinetics of the OH + Hydroxyacetone Reaction As a Function of Temperature[END_REF] hypothesized that the use of different pressures might be a source of the difference between their results at P = 2-5 Torr and those of Dillon et al. [START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF] at P= 60 Torr. Indeed, Vu et al. [START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF] observed a pronounced decrease of k 1 at total pressures lower than ∼45 Torr at T = 301 K (Figure 7) and rationalized the pressure dependence of k 1 in the frame of the reaction mechanism including the formation of a hydrogen-bonded pre-reactive complex and its collisional thermalization at high pressures. However, the results of the present study do not seem to support the pressure dependence of the reaction rate constant, and in this sense, raise more questions than give answers. Firstly, we have not observed any appreciable pressure dependence of k 1 in the pressure range 1-10.4 Torr (Figure 7). Secondly, our low pressure data (at P  2 Torr of He) for k 1 are in excellent agreement in the whole temperature range with those of Dillon et al. [START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF] and Vu et al. [START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF] measured at P = 60 and 50 Torr, respectively. In our opinion, a relatively weak pressure dependence of k 1 , if exists, can be masked by non-identified pressure independent systematic errors (up to a factor of 2 at room temperature) in some of the available studies.

A fit to the present data with modified Arrhenius expression, k = A×T n ×exp(-E a /RT), provides (blue line in Figure 8):

k 1 = 4.7×10 -22 ×T 3.25 exp(1410/T) cm 3 molecule -1 s -1 at T = 250-830K.
If our data are combined with those of Dillon et al. [START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF] and Vu et al., [START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF] one gets (black solid line in Figure 8): showing that all temperature data from these three temperature dependent measurements of k 1 fall into this range. Based on this and considering very good agreement between the k 1 data from Dillon et al., [START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF] Vu et al. [START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF] and present study, obtained under different experimental conditions and with different experimental methods, we would recommend the above expression for k 1 to use in the temperature range 230-830 K with independent of temperature 20% uncertainty. Interestingly, the temperature dependence of k 1 described by this expression is almost indistinguishable from that corresponding to the theory-based expression k 1 = 5.3×10 -20 ×T 2.6 exp(1100/T) cm 3 molecule -1 s -1 (red dashed line in Figure 8) which was proposed by Vu et al. [START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF] to represent their high-pressure k 1 data and those of Dillon et al. [START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF] over the combined range 230-500 K. It can be noted also that the results of theoretical study of Galano, [START_REF] Galano | Theoretical Study on the Reaction of Tropospheric Interest: Hydroxyacetone + OH. Mechanism and Kinetics[END_REF] reporting k 1 = 5.29×10 -23 ×T 3.4 exp(1623/T) cm 3 molecule -1 s -1 at T = 280-500K, reproduce quite well the temperature dependence of k 1 (Figure 8) observed by Dillon et al., [START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF] Vu et al. [START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF] and in the present study, although with the absolute values of k 1 approximately twice smaller.

The present study, as one can see in Figure 8, strongly supports two [START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF][START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF] of three previous temperature dependent measurements of k 1 . However the reason for the disagreement with earlier room temperature studies [START_REF] Dagaut | Kinetic Measurements of the Gas-Phase Reactions of Hydroxyl Radicals With Hydroxy Ethers, Hydroxy Ketones, and Keto Ethers[END_REF][START_REF] Orlando | The Rate and Mechanism of the Gas-Phase Oxidation of Hydroxyacetone[END_REF][START_REF] Chowdhury | ArF Laser Photodissociation Dynamics of Hydroxyacetone: LIF Observation of OH and Its Reaction Rate with the Parent[END_REF][START_REF] Butkovskaya | Mechanism of the OH-Initiated Oxidation of Hydroxyacetone over the Temperature Range 236-298 K[END_REF] and that of Baasandorj et al. [START_REF] Baasandorj | Experimental and Theoretical Studies of the Kinetics of the OH + Hydroxyacetone Reaction As a Function of Temperature[END_REF] remains unclear and even somewhat mysterious, given the very low scatter of k 1 measurements for both lower [START_REF] Dagaut | Kinetic Measurements of the Gas-Phase Reactions of Hydroxyl Radicals With Hydroxy Ethers, Hydroxy Ketones, and Keto Ethers[END_REF][START_REF] Orlando | The Rate and Mechanism of the Gas-Phase Oxidation of Hydroxyacetone[END_REF][START_REF] Chowdhury | ArF Laser Photodissociation Dynamics of Hydroxyacetone: LIF Observation of OH and Its Reaction Rate with the Parent[END_REF][START_REF] Butkovskaya | Mechanism of the OH-Initiated Oxidation of Hydroxyacetone over the Temperature Range 236-298 K[END_REF][START_REF] Baasandorj | Experimental and Theoretical Studies of the Kinetics of the OH + Hydroxyacetone Reaction As a Function of Temperature[END_REF] and higher (refs. 8,13 and present work) values of k 1 reported at T = 298 K. The reaction of OH radicals seems to be a rare case where almost all the data for the reaction rate constant were determined using the absolute measurements (except one room temperature study [START_REF] Orlando | The Rate and Mechanism of the Gas-Phase Oxidation of Hydroxyacetone[END_REF] and the present work). It seems that an in-depth study (even if only at room temperature) employing a relative rate method for the measurements of k 1 at high pressure would be helpful to verify once again the pressure dependence of k 1 and for the arbitration between the low and high values of the rate constant at T <350K.

CONCLUSIONS

In this work, we examined the kinetics of the reaction of OH radicals with hydroxyacetone using a discharge flow reactor combined with an electron impact ionization mass spectrometer. The total rate constant of the reaction, k 1 = 4.7×10 -22 ×T 3.25 exp(1410/T) cm 3 molecule -1 s -1 , was determined using both relative rate and absolute methods in the temperature range 250-830 K. The present data combined with selected previous temperature dependent studies [START_REF] Vu | Absolute Rate Coefficient of the Gas-Phase Reaction between Hydroxyl Radical (OH) and Hydroxyacetone: Investigating the Effects of Temperature and Pressure[END_REF][START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF] of the reaction (1) yield k 1 = 4.4×10 -20 ×T 2.63 exp(1110/T) cm 3 molecule -1 s - 1 , which is recommended from the present work at T = 230-830K (with conservative uncertainty of 20% at all temperatures). The U-shaped temperature dependence of k 1 observed in the present work at 2 Torr total pressure of He is in excellent agreement with two previous studies 8, [START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF] of reaction (1) at P = 50-60 Torr in contrast to the positive temperature dependence observed by Baasandorj et al. [START_REF] Baasandorj | Experimental and Theoretical Studies of the Kinetics of the OH + Hydroxyacetone Reaction As a Function of Temperature[END_REF] at T = 280-350 K and P = 2-5 Torr. In addition, the present measurements provide by a factor of 2 higher value of k 1 at T = 298K compared with those reported in earlier room temperature studies. [START_REF] Dagaut | Kinetic Measurements of the Gas-Phase Reactions of Hydroxyl Radicals With Hydroxy Ethers, Hydroxy Ketones, and Keto Ethers[END_REF][START_REF] Orlando | The Rate and Mechanism of the Gas-Phase Oxidation of Hydroxyacetone[END_REF][START_REF] Chowdhury | ArF Laser Photodissociation Dynamics of Hydroxyacetone: LIF Observation of OH and Its Reaction Rate with the Parent[END_REF][START_REF] Butkovskaya | Mechanism of the OH-Initiated Oxidation of Hydroxyacetone over the Temperature Range 236-298 K[END_REF] The reasons for these disagreements remain unclear. The present results do not support the hypothesis that the divergence between low and high pressure temperature dependent studies is due to dependence of k 1 on pressure. 8

ASSOCIATED CONTENT

Supporting Information

Diagram of the high temperature flow reactors (Figure S1); Kinetics of OH consumption at T = 330 K (Figure S2); Second-order plots for the OH + hydroxyacetone reaction at T = 298,

[Br 2 ]

 2 upon OH conversion to HOBr in reaction 3: [OH] = [HOBr] = [Br 2 ]. The method employed for OH detection and their absolute calibration was discussed in details in previous

3. 1 . 5 - 3 . 5 )

 1535 Rate constant of reaction (1): absolute measurements. The rate constant of the title reaction was determined at total pressure of (2.0-2.1) Torr under pseudo-first order conditions monitoring OH decays in excess of hydroxyacetone. Temporal profiles of OH (Figures 1 and S2 (SI)) were fitted to an exponential function [OH] = [OH] 0 ×exp(-k 1 '×t), where [OH] 0 and [OH] are the initial and time-dependent concentrations of OH radicals, respectively, and k 1 ' = k 1 ×[HAC]+k w is the pseudo-first-order rate constant with k w representing the wall loss of OH radicals, which could be measured in the absence of HAC in the reactor. In these experiments, the initial concentration of OH radicals was in the range (1.×10 11 molecule cm -3 ; the ranges of hydroxyacetone concentrations are shown in Table

Figure 1 .

 1 Figure 1. Examples of the kinetics of OH consumption in reaction with hydroxyacetone: T = 500 K. The flow velocity in the reactor was in the range (1190-2190) cm s -1 . The consumption of hydroxyacetone in reaction (1) was insignificant due to its sufficiently high excess over OH radicals. The probable thermal decomposition of HAC at high temperatures of the study was also explored. Firstly, we have not observed any visible consumption of hydroxyacetone on the time scale of our experiments. The signals at m/z = 94/96, 96/98, 110/112 and 122/124, corresponding to CH 3 Br, HOBr, CH 2 BrOH and CH 3 C(O)Br, respectively, which could be formed in reactions of possible products of HAC decomposition in presence of Br 2 , were also analyzed. The concentrations of these species were estimated to be lower than 10 11 molecule cm -3 at highest concentrations of HAC and highest temperatures of the study. Based on these observations, the thermal decomposition of hydroxyacetone could be considered as insignificant under experimental conditions of the study and not affecting the kinetic measurements. Examples of the second-order plots at different temperatures are shown in Figures 2 and

Figure 2 .

 2 Figure 2. Dependence of the pseudo-first-order rate constant (k 1 ′) on the concentration of hydroxyacetone at different temperatures. Error bars correspond to typical uncertainties (≤ 5%) on the determination of k 1 ′. For clarity, k 1 ′ data at T = 630 and 500K are Y-shifted by -20 and -45 s -1 , respectively.

  [OH] 0 = 1.4 10 11 molecule cm -3 500 K; [OH] 0 = 3.2 10 11 molecule cm -3

( 1 )

 1 : relative measurements. Considering the potential heterogeneous complications in the absolute measurements of the reaction rate constant, we have conducted two series of the relative rate measurements with reactions of OH with Br 2 and OH + n-heptane as a reference. In the experiments with OH + Br 2 as a reference reaction, the initial concentration of OH radicals, [OH] 0 , was titrated by a mixture of Br 2 and HAC and the concentration of HOBr was measured as a function of the [HAC]/[Br 2 ] ratio: OH + CH 3 C(O)CH 2 OH  products HOBr formed corresponds to the fraction of [OH] 0 consumed in reaction with Br 2 : HOBr HAC OH Rearrangement of this expression leads to: keep the contribution of the second term in equation (I) constant, the experiments were carried out at a constant concentration of Br 2 . k 1 /k 3 ratios were derived from the plots of ([OH] 0 /[HOBr] -1) as a function of the [HAC]/[Br 2 ] ratio. In the experiments, HOBr was monitored in both HAC -free system, corresponding to [OH] 0 , and in the Br 2 and HACcontaining system, corresponding to the fraction of [OH] 0 reacted with Br 2 . The measurements were carried out at a total pressure of nearly 2 Torr of He in the reactor. Initial concentration of OH radicals was in the range (2-6)×10 11 molecule cm -3 , reaction time was (0.02 -0.05) s, concentrations of Br 2 and hydroxyacetone are shown in Table 2. Examples of the experimental data are displayed in Figures 3 and S4 (SI). The values of k 1 /k 3 at different temperatures were derived from the slopes of the linear dependences in Figures 3 and S4, according to equation (I). Final values of k 1 (Table

Figure 3 .

 3 Figure 3. HOBr formed upon OH consumption by Br 2 + HAC mixtures at different temperatures. The error bars represent typical uncertainties of the measurements (≤ 10%).

  ) are expected to be very slow at the low temperatures of the study. The observed independence of the experimental results of the reaction time, which was varied from 0.015 to 0.06 s, seems to support this assumption. It can also be noted that the possible secondary chemistry initiated by oxygen atoms formed in the self-reaction of OH radicals,OH + OH  O + H 2 O (5)has a limited impact on the concentrations of HAC and n-heptane, because O-atoms are mainly consumed in the fast reactions with OH radicals and NO 2 present in high concentrations in the reactor:O + NO 2  NO + O 2(6)k 6 = 9.8510 -16 T 1.41 exp(543/T) cm 3 molecule -1 s -1 (T = 220 -950 K),[START_REF] Bedjanian | Rate Constants of the Reactions of O( 3 P) Atoms with Br 2 and NO 2 over the Temperature Range 220-950 K[END_REF] O + OH  H + O 2(7) k 7 = 1.810 -11 exp(180/T) cm 3 molecule -1 s -1 (T = 136 -515 K),[START_REF] Burkholder | Chemical Kinetics and Photochemical Data for Use in Atmospheric Studies[END_REF] H + NO 2  OH + NO(2)The observed dependences of ln([HAC] 0 /[HAC]) on ln([n-heptane] 0 /[n-heptane]) are presented in Figure4. According to equation (II) the slopes of the straight lines in Figure4correspond to the k 1 to k 4 ratio. For the rate constant of the reference reaction (4) we have used k 4 = 3.7510 -16 T 1.65 exp(101/T) cm 3 molecule -1 s -1 recommended at T = 220 -1300 K

Figure 4 .

 4 Figure 4. Dependence of ln([HAC] 0 /[HAC]) on ln(n-heptane] 0 /[n-heptane]) observed at P = 2.1 Torr and T = 290 and 320 K.

3. 3 .

 3 Rate constant of reaction (1): pressure dependence. The rate constant of reaction (1) was measured at total pressure of helium between 1.0 and 10.4 Torr and T = 301 K. This temperature was chosen for comparison with data of Vu et al. 8 who reported the measurements of k 1 at this temperature and pressures between 10 and 70 Torr. Relative rate methods, as described above, were applied in these measurements. The experimental data observed with the reaction of OH with Br 2 as a reference are shown in Figure 5. One can note that the results obtained at different pressures are similar in the range of the experimental uncertainty. This results in very close values for k 1 at different pressures: 6.06, 6.02, 6.44 and 5.8410 -12 cm 3 molecule -1 s -1 at P = 1, 2, 5 and 10.4 Torr of He, respectively, with estimated overall uncertainty on k 1 of 20%.

Figure 5 .

 5 Figure 5. HOBr yield from OH radical titration with Br 2 + HAC mixtures at T = 301 K and different pressures. The error bars represent typical uncertainties of the measurements (≤ 10%).

1 s - 1

 1 at P = 1.45, 3.25 and 7.6 Torr of He, respectively.

Figure 6 .

 6 Figure 6. Dependence of ln([HAC] 0 /[HAC]) on ln(n-heptane] 0 /[n-heptane]) observed at different pressures and T = 301 K. Partially shown error bars correspond to maximum uncertainty of 5% on the measurements.

Figure 7 . 3 . 4 .

 734 Figure 7. Reaction OH + CH 3 C(O)CH 2 OH: pressure dependence of the reaction rate constant at T = 301K.

Figure 8 .

 8 Figure 8. Reaction OH + CH 3 C(O)CH 2 OH: temperature dependence of the reaction rate constant. AM: absolute measurement, RM: relative rate method. The error bars for the present data correspond to estimated 15 and 20 % uncertainty on the measurements of k 1 by absolute and relative rate method, respectively; uncertainties on previous measurements are given as reported in respective studies. Lines correspond to the fit with modified Arrhenius expression, k = A×T n ×exp(-E a /RT), of the data from the present study (solid blue line), the present study combined with those of Dillon et al. 13 and Vu et al. 8 (black solid line) with 20% uncertainty (dashed black lines) and to theory-based expression proposed by Vu et al. 8 (red dashed line).

  et al.[START_REF] Dillon | Reaction of HO with Hydroxyacetone (HOCH 2 C(O)CH 3 ): Rate Coefficients (233-363 K) and Mechanism[END_REF] / 233-363 Baasandorj et al. 14 / 280-350 Vu et al. 8 / 293-498 This work (AM) / 298-830 This work (RM/Br 2 ) / 250-750 This work (RM/n-heptane) / 290-320 Galano 15 / 280-500 / Theory

k 1 = 4 .

 14 4×10 -20 ×T 2.63 exp(1110/T) cm 3 molecule -1 s -1 at T = 230-830K.The dashed black lines in Figure8correspond to ± 20% deviation from this expression,

Table 1 . Summary of the Absolute Measurements of the Rate Constants of the Reaction OH + Hydroxyacetone.

 1

	T (K)	Number of kinetic runs	[HAC] (10 13 molecule cm -3 )	a k 1 (10 -12 cm 3 molecule -1 s -1 )
	298	12	0.40-6.64	6.03
	330	10	0.23-5.69	5.18
	336	10	0.25-5.68	5.32
	370	10	0.17-5.66	4.73
	425	11	0.20-4.21	4.41
	500	16	0.17-6.94	4.64
	565	11	0.13-6.18	5.20
	630	11	0.15-5.08	5.76
	715	9	0.32-4.85	6.90
	830	9	0.15-5.09	7.59

a Statistical uncertainty on k 1 is ≤ 5% and the total estimated uncertainty is approximately 15%.

Table 2 . Summary of the Relative Measurements of the Rate Constant of the Reaction OH + hydroxyacetone with reaction OH + Br 2 as a reference.

 2

	d

T (K) a No./exp. b [HAC] c [Br 2 ] c k 1 /k 3 k 1 a T = 250-301 K: halocarbon wax coated reactor; T = 310-750 K: uncoated quartz reactor. b Number of data points. c Units of 10 13 molecule cm -3 . d Units of 10 -12 cm 3 molecule -1 s -1 , estimated uncertainty on k 1 is nearly 20%.

and 370 K (FigureS3); HOBr yield from OH radical titration with Br 2 + hydroxyacetone mixtures at T = 165, 465 and 555 K (FigureS4).

ACKNOWLEDGEMENT

This work was partly supported by the French national program LEFE/INSU.


Yuri Bedjanian
email: yuri.bedjanian@cnrs-orleans.fr

Temperature Dependent Rate Constants for the Reactions of Chlorine Atom with Methanol and Br 2

Keywords: Cl, CH 3 OH, Br 2, kinetics, rate coefficient

Kinetics of the reaction of Cl atoms with methanol has been investigated at 2 Torr total pressure of helium and over a wide temperature range 225 -950 K, using a discharge flow reactor combined with an electron impact ionization quadrupole mass spectrometer. The rate constant of the reaction Cl + CH 3 OH  products (1) was determined using both absolute measurements under pseudo-first order conditions, monitoring the kinetics of Clatom consumption in excess of methanol and relative rate method, k 1 = (5.1±0.8)10 -11 cm 3 molecule -1 s -1 , and was found to be temperature independent over the range T = 225-950K.

The rate constant of the reaction Cl + Br 2  BrCl + Br (3) was measured in an absolute way monitoring Cl-atom decays in excess of Br 2 : k 3 = 1.6410 -10 exp(34/T) cm 3 molecule -1 s -1 at T = 225 -960 K (with conservative 15% uncertainty). The experimental data for k 3 can also be adequately represented by the temperature independent value of k 3 = (1.8±0.3) × 10 -10 cm 3 molecule -1 s -1 . The kinetic data from the present study are compared with previous measurements.

INTRODUCTION

Methanol (CH 3 OH) is a smallest representative of alcohols, chemical species belonging to a large family of Oxygenated Volatile Organic Compounds playing an important role in the tropospheric chemistry. [START_REF] Calvert | Mechanisms of atmospheric oxidation of the oxygenates[END_REF] CH 3 OH is one of the most abundant volatile organic species in the atmosphere due to important biogenic emissions and in situ production in atmospheric oxidation of methane and other organic compounds. [START_REF] Jacob | Global budget of methanol: Constraints from atmospheric observations[END_REF] Atmospheric degradation of methanol is triggered mainly by its reaction with OH radicals, [START_REF] Calvert | Mechanisms of atmospheric oxidation of the oxygenates[END_REF] however contribution of the reaction with Cl atoms can be significant in certain locations with elevated concentration of Cl. [START_REF] Spicer | Unexpectedly high concentrations of molecular chlorine in coastal air[END_REF][START_REF] Ravishankara | Are chlorine atoms significant tropospheric free radicals?[END_REF][START_REF] Raff | Chlorine activation indoors and outdoors via surfacemediated reactions of nitrogen oxides with hydrogen chloride[END_REF] Hence, the information on kinetics and products of Cl reaction with methanol in an extended temperature range is of interest also for laboratory studies of other reactions important in atmospheric and combustion chemistry.

Kinetics of the reaction

Cl + CH 3 OH  products (1) has been extensively studied using both absolute and relative rate methods. [START_REF] Michael | Rate constants for the reaction of atomic chlorine with methanol and dimethyl ether from 200 to 500 K[END_REF][START_REF] Payne | Kinetics of the reactions of atomic chlorine with methanol and the hydroxymethyl radical with molecular oxygen at 298 K[END_REF][START_REF] Wallington | Gas phase reaction of Cl atoms with a series of oxygenated organic species at 295 K[END_REF][START_REF] Lightfoot | Flash photolysis study of the methylperoxy + hydroperoxy reaction between 248 and 573 K[END_REF][START_REF] Nelson | Absolute and relative rate constants for the reactions of hydroxyl radicals and chlorine atoms with a series of aliphatic alcohols and ethers at 298 K[END_REF][START_REF] Dóbé | Fast flow kinetic studies of the reaction CH 2 OH + HCl ⇋ CH 3 OH + Cl. The heat of formation of hydroxymethyl[END_REF][START_REF] Tyndall | Rate coefficients for the reactions of chlorine atoms with methanol and acetaldehyde[END_REF][START_REF] Smith | Infrared laser absorption measurements of HCl(v=1) production in reactions of Cl atoms with isobutane, methanol, acetaldehyde, and toluene at 295 K[END_REF][START_REF] Seakins | Rate coefficients and production of vibrationally excited HCl from the reactions of chlorine atoms with methanol, ethanol, acetaldehyde and formaldehyde[END_REF][START_REF] Taketani | Kinetics of the reactions of Cl*( 2 P 1/2 ) and Cl( 2 P 3/2 ) atoms with CH 3 OH, C 2 H 5 OH, n-C 3 H 7 OH, and i-C 3 H 7 OH at 295 K[END_REF][START_REF] Garzón | Atmospheric reactions Cl+CH 3 -(CH 2 ) n -OH (n=0-4): A kinetic and theoretical study[END_REF][START_REF] Kaiser | Rate constant of the reaction of chlorine atoms with methanol over the temperature range 291-475 K[END_REF][START_REF] Hui | Temperature dependence of the reaction of chlorine atoms with CH 3 OH and CH 3 CHO[END_REF] Room temperature data for the reaction rate constant from different studies are consistent with each other; [START_REF] Payne | Kinetics of the reactions of atomic chlorine with methanol and the hydroxymethyl radical with molecular oxygen at 298 K[END_REF][START_REF] Wallington | Gas phase reaction of Cl atoms with a series of oxygenated organic species at 295 K[END_REF][START_REF] Lightfoot | Flash photolysis study of the methylperoxy + hydroperoxy reaction between 248 and 573 K[END_REF][START_REF] Nelson | Absolute and relative rate constants for the reactions of hydroxyl radicals and chlorine atoms with a series of aliphatic alcohols and ethers at 298 K[END_REF][START_REF] Dóbé | Fast flow kinetic studies of the reaction CH 2 OH + HCl ⇋ CH 3 OH + Cl. The heat of formation of hydroxymethyl[END_REF][START_REF] Tyndall | Rate coefficients for the reactions of chlorine atoms with methanol and acetaldehyde[END_REF][START_REF] Smith | Infrared laser absorption measurements of HCl(v=1) production in reactions of Cl atoms with isobutane, methanol, acetaldehyde, and toluene at 295 K[END_REF][START_REF] Seakins | Rate coefficients and production of vibrationally excited HCl from the reactions of chlorine atoms with methanol, ethanol, acetaldehyde and formaldehyde[END_REF][START_REF] Taketani | Kinetics of the reactions of Cl*( 2 P 1/2 ) and Cl( 2 P 3/2 ) atoms with CH 3 OH, C 2 H 5 OH, n-C 3 H 7 OH, and i-C 3 H 7 OH at 295 K[END_REF] however the temperature dependence of the rate constant is not well established. [START_REF] Michael | Rate constants for the reaction of atomic chlorine with methanol and dimethyl ether from 200 to 500 K[END_REF][START_REF] Garzón | Atmospheric reactions Cl+CH 3 -(CH 2 ) n -OH (n=0-4): A kinetic and theoretical study[END_REF][START_REF] Kaiser | Rate constant of the reaction of chlorine atoms with methanol over the temperature range 291-475 K[END_REF][START_REF] Hui | Temperature dependence of the reaction of chlorine atoms with CH 3 OH and CH 3 CHO[END_REF] As noted in a most recent temperature dependent study of reaction (1),

"additional experiments that cover temperatures both below and above room temperature are needed for further assessment." [START_REF] Hui | Temperature dependence of the reaction of chlorine atoms with CH 3 OH and CH 3 CHO[END_REF] In this regard, the objective of the present work was to measure the rate constant of reaction (1) in a wide temperature range in order to harmonize existing experimental data on temperature dependence of k 1 .

In addition, the rate constant of the reaction of Cl atoms with Br 2 was measured as a part of this study. In the present work, this reaction was used for indirect detection of Cl by mass spectrometry via conversion of Cl-atom to BrCl and as a reference reaction in the relative measurements of k 1 :

Cl + Br 2  BrCl + Br (3) Previously, the reaction rate constant was measured only at relatively low temperatures (T < 400K): the available data for k 3 are spread within a factor of 1.4. [START_REF] Clyne | Atomic resonance fluorescence spectrometry for the rate constants of rapid bimolecular reactions. Part 2.-Reactions Cl + BrCl, Cl + Br 2 , Cl + ICl, Br + IBr, Br + ICl[END_REF][START_REF] Bemand | Atomic resonance fluorescence spectrometry for rate constants of rapid bimolecular reactions[END_REF][START_REF] Nicovich | Kinetics of the reactions of O( 3 P) and Cl( 2 P) with HBr and Br 2[END_REF][START_REF] Bedjanian | Low-pressure study of the reaction of Cl atoms with isoprene[END_REF][START_REF] Khamaganov | Rate coefficients for the reactions CH 3 + Br 2 (224-358 K), CH 3 CO + Br 2 (228 and 298 K), and Cl + Br 2 (228 and 298 K)[END_REF] In the present work, we report the measurements of the rate constant in an extended temperature range, T = 225 -960 K.

EXPERIMENTAL

Experiments have been conducted at total pressure of 2 Torr of Helium using a conventional fast-flow reactor combined with a molecular beam sampling electron-impact ionization mass spectrometer. [START_REF] Bedjanian | Low-pressure study of the reaction of Cl atoms with isoprene[END_REF][START_REF] Riffault | Kinetics and mechanism of the reaction of Cl atoms with HO 2 radicals[END_REF][START_REF] Riffault | Kinetic and mechanistic study of the X and XO (X = Cl, Br) reactions with dimethyl sulfoxide[END_REF][START_REF] Morin | Experimental study of the reactions of OH radicals with propane, n-pentane, and n-heptane over a wide temperature range[END_REF] Depending on the temperature range two different flow reactors were used. Low temperature reactor, used at T = 225 -320 K, consisted of a Pyrex tube (45 cm length and 2.4 cm i.d.) surrounded by a jacket through which a thermostated ethanol was circulated (Figure S1, Supporting Information (SI)). To reduce the heterogeneous loss of Cl, the inner surface of the reactor as well as of the movable injector of Cl atoms was coated with halocarbon wax. High temperature reactor, used at T = 335 -960 K, consisted of a Quartz tube (45 cm length and 2.5 cm i.d.) with an electrical heater and water-cooled attachments (Figure S2, SI). [START_REF] Morin | Experimental study of the reactions of OH radicals with propane, n-pentane, and n-heptane over a wide temperature range[END_REF] Chlorine atoms were generated by dissociation of Cl 2 , diluted in He, in a microwave discharge and introduced into the main reactor through the movable injector as shown in This method of Cl-atom detection is free of the possible complications which may arise from the fragmentation of the chlorinated species in the ion source. It can also be noted that the concentration of BrCl is not affected by side reactions, for example, with organic radicals formed in reaction (1), since they are rapidly removed by Br 2 , which is present in the reaction system in high concentrations. Finally, Cl atoms were detected also at m/z = 162/164 (ICl + ) upon their titration with an excess of I 2 ([I 2 ] = (1.0-1.5)×10

' = k 1 ×[CH 3 OH]+k w ,
where k w represents the loss of Cl-atom in the absence of methanol in the reactor. Figure 1 displays examples of the exponential decays of Cl-atoms in the presence of different concentrations of CH 3 OH in the reactor. Examples of second-order plots observed at different temperatures are shown in Figure 2. All the pseudo-first order rate constants, k 1 ', were corrected for axial and radial diffusion of Cl atoms [START_REF] Kaufman | Kinetics of elementary radical reactions in the gas phase[END_REF] with diffusion coefficient of Cl in He calculated as D 0 = 504 × (T/298) 1.85 Torr cm -2 s -1 . [START_REF] Tang | Compilation and evaluation of gas phase diffusion coefficients of reactive trace gases in the atmosphere: volume 1. Inorganic compounds[END_REF] The corrections were typically within a few percent, being somewhat higher in a few kinetic runs (up to 15%). The results of the absolute measurements of k 1 are shown in Table 1. The measurements of k 1 at T ≥ 360 K have been conducted in an uncoated quartz reactor and those between 225 and 320 K in a Pyrex reactor coated with halocarbon wax. The combined uncertainty on k 1 was estimated to be about 15% by adding in quadrature statistical error (≤ 5%) and those on the measurements of the absolute concentration of methanol (10%), flows (3%), pressure (2%), temperature (1%).

The possible impact of secondary chemistry on the measurements of k 1 can be discussed. Firstly, it should be noted that thermal decomposition of methanol is negligible in the temperature range of the present work. Lu et al. [START_REF] Lu | Shock tube study on the thermal decomposition of CH 3 OH[END_REF] in their shock tube study reported the rate constant of 3.2×10 -9 exp(-29470/T) cm 3 molecule -1 s -1 for the thermal decomposition of CH 3 OH at T = 1660-2050 K. Extrapolation of this expression to 950 K, the highest temperature of the present study, provides the rate constant of 10 -22 cm 3 molecule -1 s -1 , which is too low to affect the measurements of k 1 . Secondary reactions involving CH 2 OH radical formed in reaction (1a) can lead both to production and consumption of Cl atoms in reactions of CH 2 OH with Cl 2 (precursor of Cl) and Cl, respectively:

Cl + CH 3 OH  CH 2 OH + HCl (1a) CH 2 OH + Cl 2  CH 2 ClOH + Cl (5) CH 2 OH + Cl  CH 2 O + HCl (6)
The possible reproduction of Cl atoms in reaction ( 5 Among the data presented in Table 1, those at T = 360 and 770 K were obtained in the presence of O 2 in the reactor.

To our knowledge the rate constant of reaction (6) was reported in two studies at room temperature: k 6 = 1.410 -10 32 and 6.6410 -10 cm 3 molecule -1 s -1 . [START_REF] Pagsberg | UV spectrum and kinetics of hydroxymethyl radicals[END_REF] The last, rather high, value of k 6 was extracted from the simulation of the experimental profiles of CH 2 OH observed upon the pulse radiolysis of SF 6 /HCI/CH 3 OH mixtures. [START_REF] Pagsberg | UV spectrum and kinetics of hydroxymethyl radicals[END_REF] With k 6 as high as 6.610 -10 cm 3 molecule -1 s -1 , in the present experiments, the contribution of the secondary reaction of Cl with CH 2 OH to the measured Cl decays would be significant, especially at low concentrations of methanol. This would lead to increase of the intercepts in the second-order plots shown in Figure 2. In fact, this was not the case: the intercepts of the second-order plots were always in good agreement with the values of k w measured in the absence of CH 3 OH in the reactor, as shown in Figure 2. Experiments with addition of O 2 into reactive system (see above) also seem to indicate that impact of reaction ( 6) is insignificant. In the presence of O

For complete confidence in the minor role of the secondary chemistry in the absolute measurements of k 1 , we have carried out relative rate measurements under conditions where secondary reactions ( 5) and ( 6) are obviously not important.

Rate constant of reaction (1): relative measurements

In the relative measurements of k 1 , reaction of Cl atoms with Br 2 was used as a reference.

The experiments consisted in a fast consumption of Cl in reaction with a mixture of Br 2 and CH 3 OH and the measurements of BrCl yield as a function of the [CH 3 OH]/[Br 2 ] ratio:

Cl + CH 3 OH  products (1) Cl + Br 2  BrCl + Br (3)
The fraction of the initial concentration of Cl atoms, [Cl] 0 , transformed to BrCl in reaction 3, is:

BrCl CH 3 OH

Cl

After rearrangement of this expression one has: One can note that the results of the relative measurements are in excellent agreement with the absolute measurements of k 1 . It should be emphasized that secondary reactions (5) and ( 6) are not operative under the experimental conditions of the relative measurements.

Firstly, in the presence of high concentrations of Br

Although the rate constant of reaction ( 9) is unknown, it is expected to be higher than k 5 by analogy with reactions of Br 2 and Cl 2 with other organic radicals.

It can also be noted that reaction of CH 3 OH with Br atoms formed in reactions (1) and ( 9): [START_REF] Dóbé | Direct kinetic studies of the reactions Br + CH 3 OH and CH 2 OH + HBr: The heat of formation of CH 2 OH[END_REF] Br + CH 3 OH  CH 2 OH + HBr (10) k 10 = 5.6610 -15 (T) 1.5 exp(-3600/T) cm 3 molecule -1 s -1 (T = 439 -713 K) is relatively slow, and, in any case, has a negligible impact on the concentrations of CH 3 OH, considering that [CH 3 OH] >> [Br] [Cl] 0 .

Rate constant of reaction (3)

The rate constant of reaction (3) was determined from Cl decays ([Cl] 0 = (0.5 -1.5)×10 [START_REF] Dóbé | Fast flow kinetic studies of the reaction CH 2 OH + HCl ⇋ CH 3 OH + Cl. The heat of formation of hydroxymethyl[END_REF] molecule cm -3 ) monitored in excess of Br 2 (Table 3). Examples of the kinetics of Cl consumption observed at different concentrations of Br 2 are shown in Figure S4 (SI). The consumption of Br 2 was generally within a few percent: only in a few measurements (at lowest concentrations of Br 2 ) it was significantly higher (up to 25 %). The reaction rate constant was always calculated using the mean concentration of Br 2 along the reaction zone.

Typical plots of the pseudo-first-order rate constant, k 3 ' = k 3 ×[Br 2 ]+k w , versus concentration of Br 2 are shown in Figure 4. Diffusion corrections on k 3 ' were usually less than 10% and up to 16% in a few kinetic runs. The Y-intercepts of all the second-order plots, k w , were in the range (0 -22) s -1 , in good agreement with the rate of Cl decay measured in the absence of Br 2 in the reactor. All the results of the absolute measurements of k 3 (with estimated total uncertainty on k 3 of 15%) are shown in Table 3. At T ≤ 320 and ≥ 340 K, the measurements were carried out in a Pyrex reactor coated with halocarbon wax and in an uncoated quartz reactor, respectively. 20 were reported at T = 298 K. The difference between two values measured with the same method and in the same group was explained [START_REF] Bemand | Atomic resonance fluorescence spectrometry for rate constants of rapid bimolecular reactions[END_REF] by Br 2 loss in the flowmeter when passing a dilute mixture of Br 2 with inert gas, leading to an underestimation of k 3 in the study of Clyne and Cruse. [START_REF] Clyne | Atomic resonance fluorescence spectrometry for the rate constants of rapid bimolecular reactions. Part 2.-Reactions Cl + BrCl, Cl + Br 2 , Cl + ICl, Br + IBr, Br + ICl[END_REF] Nicovich and Wine 21 employed a laser flash photolysis-resonance fluorescence technique to measure the rate constant of reaction (3) as a function of temperature at T = 298-401 K. k 3 = (1.58 ± 0.22)10 -10 cm 3 molecule -1 s -1 was found to be virtually independent of temperature, although a slight positive temperature dependence was observed (E/R = 144 K). [START_REF] Nicovich | Kinetics of the reactions of O( 3 P) and Cl( 2 P) with HBr and Br 2[END_REF] The k 3 data of Nicovich and Wine 21 shown in Figure 5 have been adjusted upward by 11%, as proposed by Khamaganov and Crowley, [START_REF] Khamaganov | Rate coefficients for the reactions CH 3 + Br 2 (224-358 K), CH 3 CO + Br 2 (228 and 298 K), and Cl + Br 2 (228 and 298 K)[END_REF] considering currently recommended [START_REF] Burkholder | Chemical kinetics and photochemical data for use in atmospheric studies[END_REF] absorption cross section for Br 2 at 415.8 nm of 6.5 × 10 As one can see in Figure 5, the values of k 3 from the earlier study [START_REF] Bedjanian | Low-pressure study of the reaction of Cl atoms with isoprene[END_REF] are by up to 30% lower than those measured in the present work. The reason for this disagreement is most likely due to a systematic error in previous measurements, which is difficult to identify at present. The dashed lines in Figure 5 correspond to ± 15% deviation from this expression, showing that practically all temperature dependent data for k 3 (except earlier measurements from this group) [START_REF] Bedjanian | Low-pressure study of the reaction of Cl atoms with isoprene[END_REF] fall into this range. Based on this, we would recommend the above expression for k 3 for use in the temperature range 225-960 K with conservative 15% uncertainty. The observed small negative temperature dependence of k 3 may be considered as an indication of a complex reaction mechanism that includes formation of an intermediate energized complex followed either by its dissociation back to reactants or to reaction products.

Comparison with previous studies

The results of room temperature measurements of k The reaction rate constant was found to be k 1 (±1) = (6.33±0.70)10 -11 cm 3 molecule -1 s -1 independent of temperature over the range 200-500 K. In contrast, Garzón et al. [START_REF] Garzón | Atmospheric reactions Cl+CH 3 -(CH 2 ) n -OH (n=0-4): A kinetic and theoretical study[END_REF] in their laser photolysis-resonance fluorescence study of reaction (1) reported significant temperature dependence for k 1 at T = 266-380 K: k 1 = (3.55±0.22)10 -10 exp(-559±40)/T) cm 3 molecule -1 s -1 . Kaiser and Wallington [START_REF] Kaiser | Rate constant of the reaction of chlorine atoms with methanol over the temperature range 291-475 K[END_REF] employed relative rate technique with C 2 H 6 as the reference compound and reported a weaker temperature dependence over the temperature range 291-475 K: k 1 = (8.6±1.3)10 -11 exp(-167±60)/T) cm 3 molecule -1 s -1 . Finally in a most recent study of Hui et al., [START_REF] Hui | Temperature dependence of the reaction of chlorine atoms with CH 3 OH and CH 3 CHO[END_REF] the rate constant of reaction (1) was determined from the kinetics of HO 2 formation upon pulsed laser photolysis of Cl 2 in presence of CH 3 OH and O 2 : k 1 = (5.02±1.5)10 -11 exp(20±88)/T) cm 3 molecule -1 s -1 over the range T = 230-297 K. Within experimental uncertainty, the rate constant was temperature independent with the average value of k 1 = (5.45±0.37)10 -11 cm 3 molecule -1 s -1 over this temperature range.

As one can see in Figure 6, on the absolute basis the k 1 data from the present work are in agreement within stated experimental uncertainties with all the previous temperature dependent measurements, except a few points from Garzón et al. [START_REF] Garzón | Atmospheric reactions Cl+CH 3 -(CH 2 ) n -OH (n=0-4): A kinetic and theoretical study[END_REF] The present data obtained in a wide temperature range seem to definitely confirm the temperature independence of k 1 reported in previous studies of Michael et al. [START_REF] Michael | Rate constants for the reaction of atomic chlorine with methanol and dimethyl ether from 200 to 500 K[END_REF] and Hui et al. [START_REF] Hui | Temperature dependence of the reaction of chlorine atoms with CH 3 OH and CH 3 CHO[END_REF] and question quite a strong temperature dependence of k 1 measured by Garzón et al. [START_REF] Garzón | Atmospheric reactions Cl+CH 3 -(CH 2 ) n -OH (n=0-4): A kinetic and theoretical study[END_REF] The observed independence of k 1 of temperature is in line with empirical correlations between the R-H bond energy and corresponding temperature dependence of the H-atom abstraction by Cl 6,7,18 and with theoretical calculations of k 1 in the range T = 300-1000 K shown in Figure 6 by dotted line. [START_REF] Jodkowski | Theoretical study of the kinetics of the hydrogen abstraction from methanol. 2. Reaction of methanol with chlorine and bromine atoms[END_REF] Solid line in Figure 6 corresponds to the mean value of k 1 from the present measurements at all temperatures which is recommended from this study with conservative 15% uncertainty (dashed lines): k 1 = (5.1±0.8)10 -11 cm 3 molecule -1 s -1 over the temperature range 225 -950 K.

There is an experimental [START_REF] Payne | Kinetics of the reactions of atomic chlorine with methanol and the hydroxymethyl radical with molecular oxygen at 298 K[END_REF][START_REF] Dóbé | Formation of methoxy and hydroxymethyl free radicals in selected elementary reactions[END_REF] and theoretical [START_REF] Jodkowski | Theoretical study of the kinetics of the hydrogen abstraction from methanol. 2. Reaction of methanol with chlorine and bromine atoms[END_REF]

CONCLUSION

In this work, we investigated the kinetics of the reaction of Cl atoms with methanol using a discharge flow reactor combined with an electron impact ionization mass spectrometer. The total rate constant of the reaction, k 1 = (5.1±0.8)10 -11 cm 3 molecule -1 s -1 , was determined using both absolute and relative rate method for temperatures in the range 225 -950 K.

Temperature dependence of k 1 , measured over a wide temperature range, seems to harmonize existing experimental data for k 1 and confirms independence of the rate constant of temperature up to 1000 K in agreement with theoretical predictions. The temperature dependence of the rate constant of the Cl + Br 2 reaction, measured as a part of this study, k 3 = 1.6410 -10 exp(34/T) cm 3 molecule -1 s -1 at T = 225 -960 K, and found to be in good agreement with most of the previous low temperature data, was extended to higher temperatures (T = 960K). uncertainties on previous measurements are given as reported in respective studies.

TABLES LEGENDS

Figuresk 3 = 1 .

 31 Figures S1 and S2 (SI). Three methods were used for the detection of chlorine atoms. Firstly, Cl atoms were detected at their parent peaks at m/z = 35/37 (Cl + ). In this case, the contribution at these peaks of Cl 2 (precursor of Cl atoms) and chlorinated products of reactions (1) and (3) due to their fragmentation in the ion source of the mass spectrometer (always operating at ca. 25 eV) should be taken into account. In another method, Cl atoms were monitored at m/z = 114/116/118 (BrCl + ) after being scavenged with an excess of Br 2 ([Br 2 ]  3×10 13 molecule cm -3 , added in the end of the reactor 5 cm upstream of the sampling cone, as shown in Figure S2 (SI)) in reaction 3: Cl + Br 2  BrCl + Br (3) k 3 = 1.6410 -10 exp(34/T) cm 3 molecule -1 s -1 (T = 225 -960 K, this work).

  Figure S3 (SI) demonstrates the second-order plot measured for reaction (1) at T = 440 K. As one can see, the results of the measurements of k 1 ' in the presence and in the absence of O 2 in the reactor are indistinguishable within the experimental error, which indicates a negligible role of reaction (5) under the experimental conditions of the measurements of k 1 .

  a constant concentration of Br 2 , the second term in equation (I) is constant and k 1 /k 3 can be determined as a slope of the linear dependence of ([Cl] 0 /[BrCl] -1) on the [CH 3 OH]/[Br 2 ] ratio. In the experiments, BrCl was monitored in both CH 3 OH -free system, corresponding to [Cl] 0 , and in the Br 2 and CH 3 OH -containing system, corresponding to the fraction of [Cl] 0 reacted with Br 2 . Reaction time was (0.012 -0.017) s, [Cl] 0 = (2-5)×10 11 molecule cm -3 , concentrations of CH 3 OH and Br 2 are shown in Table 2. The relative measurements of k 1 were carried out in an uncoated quartz reactor, except T = 227, 243 and 295 K, where the reactor coated with halocarbon wax was used. Typical examples of the experimental data are shown in Figure 3. Final values of k 1 (Table2) were calculated with k 3 = 1.6410 -10 exp(34/T) cm 3 molecule -1 s -1 determined in this study with conservative 15% uncertainty at all temperatures (T = 225 -960K).

Figure 5

 5 Figure 5 summarizes the results of the measurements of the rate constant of reaction (3). In earlier room temperature discharge-flow studies using Cl detection by atomic resonance fluorescence, the values of k 3 = (1.20 ± 0.15)10 -10 19 and (1.9 ± 0.2)10 -10 cm 3 molecule -1 s -1

Khamaganov and Crowley 23 Figure 5 )

 235 Figure5) provides the following Arrhenius expression:k 3 = 1.6410 -10 exp(34/T) cm 3 molecule -1 s -1 at T = 225 -960K.

1 7 -

 7 [START_REF] Taketani | Kinetics of the reactions of Cl*( 2 P 1/2 ) and Cl( 2 P 3/2 ) atoms with CH 3 OH, C 2 H 5 OH, n-C 3 H 7 OH, and i-C 3 H 7 OH at 295 K[END_REF] (summarized, for instance, in ref.18), employing various experimental techniques and both absolute and relative rate methods, are consistent with each other and fall in the range (4.57-6.14) )10 -11 cm 3 molecule -1 s -1 with an average value of k 1 (±2) = (5.35±0.90)10 -11 cm 3 molecule -1 s -1 at T = 298 K. The available temperature dependent k 1 data[START_REF] Michael | Rate constants for the reaction of atomic chlorine with methanol and dimethyl ether from 200 to 500 K[END_REF][START_REF] Garzón | Atmospheric reactions Cl+CH 3 -(CH 2 ) n -OH (n=0-4): A kinetic and theoretical study[END_REF][START_REF] Kaiser | Rate constant of the reaction of chlorine atoms with methanol over the temperature range 291-475 K[END_REF][START_REF] Hui | Temperature dependence of the reaction of chlorine atoms with CH 3 OH and CH 3 CHO[END_REF] are shown in Figure6along with the results of the present work. The first determination of the rate constant of reaction(1) was realized by Michael et al.[START_REF] Michael | Rate constants for the reaction of atomic chlorine with methanol and dimethyl ether from 200 to 500 K[END_REF] using the flash photolysis-resonance fluorescence technique.

FIGURE 1

 1 FIGURE 1 Reaction Cl + CH 3 OH: examples of Cl decay profiles in presence of different amounts of methanol at T = 560 K.

FIGURE 2

 2 FIGURE 2 Reaction Cl + CH 3 OH: plot of pseudo-first-order rate constant (k 1 ′) versus concentration of CH 3 OH at different temperatures. Error bars represent typical uncertainties (≤ 5%) on the determination of k 1 ′. For clarity, the data for T = 560 and 950 K are Y-shifted by 100 and 200 s -1 , respectively.

FIGURE 3

 3 FIGURE 3 Yield of BrCl from Cl-atom titration with Br 2 + CH 3 OH mixtures at different temperatures. Partially shown error bars represent typical uncertainties of the measurements (≤ 10%). For clarity, the data for T = 385 and 600 K are Y-shifted by 0.5 and 1.0, respectively.

FIGURE 4

 4 FIGURE 4 Reaction Cl + Br 2 : plot of pseudo-first-order rate constant (k 3 ′) versus concentration of Br 2 at different temperatures. Partially shown error bars represent typical uncertainties (≤ 5%) on the determination of k 3 ′.

FIGURE 5

 5 FIGURE 5 Reaction Cl + Br 2 : summary of the measurements of the temperature dependence of the reaction rate constant. Partially shown for the present data error bars correspond to estimated 15 % uncertainty on the measurements of k 3 . Solid line: Arrhenius fit to the present data; dashed lines: present fit ±15%.

FIGURE 6

 6 FIGURE 6 Reaction Cl + CH 3 OH: summary of the temperature dependent measurements of the reaction rate constant. AM: absolute measurement, RM: relative rate method. The error bars partially shown for the present data correspond to estimated 15 and 20 % uncertainty on the measurements of k 1 in the present work by absolute and relative rate method, respectively;

  Clatom reaction with CH 3 OH is a well-known and very often used in laboratory source of CH 2 OH and, in the presence of O 2 , of HO 2 radicals:

	Cl + CH 3 OH  CH 2 OH + HCl	(1a)
	Cl + CH 3 OH  CH 3 O + HCl	(1b)
	CH 2 OH + O 2  HO 2 + CH 2 O	(2)

3 RESULTS AND DISCUSSION 3.1 Rate constant of reaction (1): absolute measurements

				13 molecule cm -3 ,
	added in the end of the reactor (Figures S1 and S2, SI): 27
	Cl + I 2		ICl + I	(4)
	k 4 = (2.1 ± 0.3)×10 -10 cm 3 molecule -1 s -1 (T = 296 -365 K).
	The absolute concentrations of Cl atoms were determined using their chemical conversion to
	BrCl in reaction (3) with an excess of Br 2 : [Cl] = [BrCl] = [Br 2 ]. The absolute calibrations of
	MS signals of the stable species (CH 3 OH, Br 2 , Cl 2 and O 2 ) were derived from measured flows
	of known mixtures of the species with helium.
	The purities of the gases used were as follows: He (99.9999%, Alphagaz), passed through
	liquid nitrogen trap; O 2 (99.999%, Alphagaz); Br 2 (>99.99%, Aldrich); Cl 2 (>99%, Ucar); I 2
	(99.999%, Aldrich); methanol (> 99.9%, Aldrich).
	Experiments were carried out under pseudo-first order conditions monitoring kinetics of Cl
	consumption ([Cl] 0 = (1 -2)×10 11 molecule cm -3 ) in an excess of methanol (see Table 1 for
	the concentrations of CH 3 OH used). It was observed that generally the consumption of
	CH 3 OH was insignificant due to its sufficiently high excess over Cl atoms. Under such
	conditions, [Cl] = [Cl] 0 ×exp(-k 1 '×t), with pseudo-first-order rate constant k 1

  2 , CH 2 OH is transformed via reaction (2) into CH 2 O and HO 2 , which are less reactive toward Clatom compared with CH 2 OH (k 7 = 4.310 -11 and k 8 = 8.110 -11 exp(-30/T) cm 3 molecule -1 s -1 at

	T = 230-420 and 200-500 K, respectively): 31
	Cl + HO 2 	products	(7)
	Cl + CH 2 O 	HCO + HCl

  2 and CH 3 OH, Cl atoms are rapidly titrated in primary reactions (1) and (3). Secondly, CH 2 OH radical is expected to be rapidly scavenged by Br 2 ([Br 2 ]>>[Cl 2 ]):

	CH 2 OH + Br 2		CH 2 BrOH + Br

  -19 cm 2 molecule -1 instead of 5.87 × 10 -19 cm 2 molecule -1 used by Nicovich and Wine 21 in optical measurements of the Br 2

	concentrations. In previous study from our group, 22 essentially similar (to those of the
	present work) experimental equipment and experimental approach were used except that Cl
	atoms were titrated at the end of the reactor with C 2 H 3 Br to be monitored at m/z = 62/64
	(C 2 H 3 Cl + ):
	Cl + C 2 H 3 Br		Br + C 2 H 3 Cl	(11)

  The experiment consisted in successive titration of the same initial concentration of Cl atoms with CH 3 OH and CD 3 OH with detection of HCl formed. In the case of Cl titration with CH 3 OH, the concentration of HCl formed corresponds to the initial concentration of Cl atoms: [HCl] CH 3 OH = [Cl] 0 . HCl, if formed upon titration of Cl with CD 3 OH, should come from the H-atom abstraction from OH group, and the branching ratio for channel (1b) could be determined as the ratio [HCl] CD 3 OH /[HCl] CH 3 OH , in a relative way without measurements of the absolute concentrations. Unfortunately, the designed experiments were unsuccessful. Atboth temperatures (300 and 900K), the HCl signal monitored in Cl+CD 3 OH system was found to be relatively high and to increase with reaction time, most likely, due to isotopic exchange reactions in the gas phase and on the wall of the reactor. Additional experimental work including a thorough analysis of this phenomenon is necessary in order to measure the branching ratio of the CH 3 O forming channel of reaction(1) which is expected to be very small.

		evidence that reaction of Cl-atom with
	methanol proceeds predominantly through hydrogen atom abstraction from the methyl
	group (reaction 1a):
	Cl + CH 3 OH  CH 2 OH + HCl	(1a)
	Cl + CH 3 OH  CH 3 O + HCl	(1b)
	Dóbé et al. 36 have reported the yield of CH 3 O radical in reaction (1) of 0.01±0.01 and
	0.04±0.02 at T = 298 and 482 K, respectively. Theoretical calculations predict nearly 1% for
	the branching ratio of the CH 3 O forming channel (1b) at T = 1000 K. 35 In the present work,
	we have tried to estimate the branching ratio for channel (1b) at two temperatures, T = 300
	and 900 K.

TABLE 1

 1 Reaction Cl + CH 3 OH: summary of the absolute measurements of the rate constant

TABLE 2

 2 Reaction Cl + CH 3 OH: summary of the relative measurements of the rate constant

TABLE 3

 3 Reaction Cl + Br 2 : summary of the measurements of the rate constant TABLES

TABLE 1

 1 Reaction Cl + CH 3 OH: summary of the absolute measurements of the rate constant

	T (K)	Number of kinetic runs	[CH 3 OH] (10 12 molecule cm -3 )	a k 1 (10 -11 cm 3 molecule -1 s -1 )
	225 b	11	0.41-10.8	5.17
	245 b	7	0.78-12.6	4.96
	266 b	9	0.59-16.0	5.14
	298 b	10	1.16-17.3	5.11
	320 b	10	1.79-17.6	5.06
	360 c,d	7	1.82-13.0	5.10
	440 c	12	0.76-12.2	5.20
	495 b	9	0.42-15.5	4.92
	500 b	11	0.55-15.2	5.01
	560 c	13	0.81-11.8	5.16
	650 b	13	0.83-15.4	5.16
	770 c,d	9	1.00-13.0	5.02
	950 b	12	0.90-12.7	5.14

a Statistical uncertainty on k 1 is ≤ 5%, total estimated uncertainty is nearly 15%; b Cl atoms monitored at m/z = 116 (BrCl + ) (see text); c Cl atoms monitored at m/z = 162 (ICl + ) (see text); d In presence of O 2 in the reactor (see text).

TABLE 2

 2 Reaction Cl + CH 3 OH: summary of the relative measurements of the rate constant Units of 10 13 molecule cm -3 ; c Statistical uncertainty on k 1 /k 3 is ≤ 4%, total estimated uncertainty is nearly 15%;

	T (K)	No./exp. a	[CH 3 OH] b	[Br 2 ] b	k 1 /k 3 c	k 1 d
	227	9	0.67-8.68	1.34	0.264	5.03
	243	12	0.56-9.67	1.10	0.278	5.24
	295	11	0.65-13.6	1.22	0.273	5.02
	335	10	0.45-10.5	1.36	0.279	5.06
	355	13	0.38-8.91	1.09	0.285	5.14
	385	9	0.38-11.2	1.44	0.292	5.23
	456	11	0.40-6.77	0.62	0.282	4.98
	600	13	0.23-5.96	1.09	0.305	5.29

a Number of data points; b d Units of 10 -11 cm 3 molecule -1 s -1 , estimated uncertainty on k 1 is nearly 20% and includes that on reference reaction.

TABLE 3

 3 Reaction Cl + Br 2 : summary of the measurements of the rate constant

	T (K)	Number of kinetic runs	[Br 2 ] (10 12 molecule cm -3 )	a k 3 (10 -10 cm 3 molecule -1 s -1 )
	225 b	10	0.71-3.84	1,97
	237 b	9	0.35-3.53	1,94
	257 b	10	0.59-4.43	1,80
	280 b	9	0.69-4.41	1,74
	298 b	9	0.33-5.02	1,85
	320 b	9	0.48-4.40	1,87
	340 b	9	0.54-4.44	1,85
	370 c	7	0.46-3.78	1,71
	400 b	7	0.28-3.57	1,87
	440 b	6	0.29-3.31	1,75
	500 b	11	0.40-3.62	1,75
	567 c	11	0.21-3.66	1,74
	650 b	7	0.20-2.14	1,70
	785 b	8	0.26-2.17	1,75
	960 c	13	0.21-3.57	1,70

a Statistical uncertainty on k 3 is ≤ 5%, total estimated uncertainty is nearly 15%. b Cl atoms monitored at m/z = 35 (Cl + ) (see text); c Cl atoms monitored at m/z = 164 (ICl + ) (see text).

ACKNOWLEDGMENTS

This research did not receive any specific grant from funding agencies in the public, commercial, or not-for-profit sectors.

SUPPORTING INFORMATION

Additional supporting information may be found online in the Supporting Information section at the end of the article.