Rate Constant for the OH + CO Reaction at Low Temperatures - The

Rate constants for the reaction of OH + CO → products (1) have been measured using laser photolysis/laser-induced fluorescence (LP/LIF) over the tem...
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Rate Constant for the OH + CO Reaction at Low Temperatures Yingdi Liu and Stanley P. Sander* Jet Propulsion Laboratory, California Institute of Technology, 4800 Oak Grove Drive, Pasadena, California 91109, United States S Supporting Information *

ABSTRACT: Rate constants for the reaction of OH + CO → products (1) have been measured using laser photolysis/laser-induced fluorescence (LP/LIF) over the temperature range 193−296 K and at pressures of 50−700 Torr of Ar and N2. The reaction was studied under pseudo-first-order conditions, monitoring the decay of OH in the presence of a large excess of CO. The rate constants can be expressed as a combination of bimolecular and termolecular components. The bimolecular component was found to be temperature-independent with an expression given by kbi(T) = (1.54 ± 0.14) × 10−13[e−(13±17)/T] cm3 molecule−1 s−1, with an error of one standard deviation. The termolecular component was fitted to the expression, kter = k0(T)[M]/[1 + (k0(T)[M]/k∞(T)] × 2 −1 −n 0.6{1+[log10(k0(T)[M]/k∞(T))] } where k0(T) = k300 and k∞(T) = k300 0 (T/300) ∞ (T/ −m = (6.0±0.5) × 300) . The parameters for k0(T) were determined to be k300 0 −33 10−33 cm6 molecule−2 s−1 in N2 and k300 cm6 molecule−2 s−1 0 = (3.4 ± 0.3) × 10 in Ar, with n = 1.9±0.5 and 2.0±0.4 in N2 and Ar, respectively. These parameters were determined using k0(T) and m from the NASA kinetics data evaluation (JPL Publication No. 10-6) since the experimental pressure range was far from the high-pressure limit. Addition of low concentrations of O2 had no discernible effect on the mechanism of the OH + CO reaction but resulted in secondary reactions which regenerated OH.

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INTRODUCTION CO is an important species in Earth’s atmosphere and plays a crucial role in atmospheric chemistry.2 The reaction of CO with OH is one of the most important processes controlling the lifetime and abundance of HOx (OH + HO2) radicals in the troposphere thereby influencing the production of ozone under both polluted and unpolluted conditions. The CO + OH reaction is important in combustion chemistry where it is the principal mechanism for converting CO into CO2 in flames. This reaction is also important in the atmospheric chemistry of Mars. As shown below, CO + OH is the rate-limiting step in the cycle that controls the CO/CO2 ratio in the atmosphere of Mars3,4

pressures up to several atmospheres, the observed rate constant depends linearly on the bath gas density. The pressure dependence is explained by a mechanism involving the formation of a vibrationally excited HOCO intermediate, HOCO*, which can decompose to form H + CO2 or undergo collisional stabilization. M

(R1a)

HOCO* → H + CO2

(R1b)

In Earth’s atmosphere, the thermalized HOCO intermediate formed in channel R1a reacts rapidly with O2 to produce HO2 + CO211,14 with a bimolecular rate constant of 1.5 × 10−12 cm3 molecule−1 s−1 at ∼298 K. The H atoms formed in channel R1b react rapidly with atmospheric O2 to produce HO2. Therefore, the end products from both reaction channels are HO2 and CO2.1 Previous studies indicated that the presence of O2 has no effect on the rate constant other than as a third body.1 To establish relevance to the Martian atmosphere and Earth’s upper troposphere/lower stratosphere, rate constants are required in the temperature range 153−279 K and pressure range from a few Torr to 1 atm. However, there are few laboratory kinetics studies that cover these ranges, especially at the lowest temperatures. The lowest temperature kinetics data were obtained by Frost et al.15 who used the laser photolysis/ laser-induced fluorescence (LP/LIF) technique to measure rate

OH + CO → H + CO2 H + O2 + M → HO2 + M O + HO2 → OH + O2 Net: O + CO → CO2

The CO + OH reaction has been the subject of numerous laboratory and theoretical studies over the past several decades.1 Previous studies found that the rate constant for OH + CO at low pressure is approximately independent of temperature over the range 250−500 K, with a recommended value of 1.5 × 10−13 cm3 molecules−1 s−1 at 298 K in the limit of zero bath gas density.1 Above 500 K, the rate constant increases rapidly with temperature. An increase in the rate constant with pressure has been observed in many investigations.5−13 At © XXXX American Chemical Society

OH + CO ⇌ HOCO* → HOCO

Received: July 25, 2015 Revised: August 13, 2015

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The Journal of Physical Chemistry A Table 1. CO + OH Reaction Rate Constant Summarya kbi ± σ

bath gas 1

a

NASA this work

air N2 Ar

IUPAC18 McCabe et al.2

air or N2 air or N2

−13

(1.5 ± 0.2) × 10 × (T/300) (1.54 ± 0.14) × 10−13

k0300 ± σ (−0.6±0.1)

n±σ −33

(5.9 ± 0.6) × 10 1.4 ± 0.1 (6.0 ± 0.5) × 10−33 1.9 ± 0.5 (3.4 ± 0.3) × 10−33 2.0 ± 0.4 k 1.44 × 10−13(1 + [M]/4.2 × 1019) 1.57 × 10−13(1 + [M] /4.4 × 1019)

k∞300 ± σ (1.1 ± 0.1) × 10−12 (1.1 ± 0.1) × 10−12 b (1.1 ± 0.1) × 10−12 b

m±σ −1.3 ± 0.1

Constants units: first order, s−1; second order, cm3 molecule−1 s−1; third order, cm6 molecule−2 s−1. bValue adopted from ref 1.

constants for the OH + CO and OD + CO reactions over the temperature range 80−297 K (for OH) at a total pressure of 2−10 Torr of N2 and Ar. The absolute accuracy of the rate constant measurements from this study at low temperatures was estimated to be ±50%. Fulle et al. studied OH + CO between 90 and 830 K in He bath gas over the pressure range 1−700 bar.16 Senosiain et al.17 carried out master equation calculations on a theoretical potential energy surface that reproduced most of the experimental data, permitting extrapolation of the rate coefficient outside the ranges of experimental pressure and temperature. Although their model closely reproduces many experiments conducted at room temperature and above, it underpredicts rate constants below 250 K. The study by Senosiain et al. forms the basis for the rate constant expressions recommended in the NASA data evaluation.1 The rate constants recommended by IUPAC for atmospheric conditions18 include a pressure-dependent term but are temperature-independent over the range 200−300 K, in contrast to the NASA evaluation, which recommends a weak temperature dependence (as shown in Table 1). In the present study, rate constants were measured for the reaction of CO with OH using the laser photolysis/laser-induced fluorescence (LP/LIF) technique between 193−293 K over the pressure range 50− 700 Torr in Ar and N2. Temperature dependent falloff parameters were derived. In addition, the effect of O2 on the observed OH kinetics was analyzed.

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or H 2O2 → 2OH (248 nm, φ ≈ 1.7)

Anhydrous HNO3 was introduced by passing carrier gas (N2 or Ar) through a bubbler containing a mixture of HNO3 (70 wt % Sigma-Aldrich).) and H2SO4 (98 wt % Sigma-Aldrich). The H2SO4 (98 wt %)/HNO3 (70 wt %) vol/vol ratio was 3:1, which ensured the complete removal of water from the HNO3 solution. The initial OH concentration, [OH]0 from the photolysis of HNO3 was (3−6) × 1010 molecules cm−3 on the basis of the photolysis laser fluence and HNO3 concentration calculated from its saturation vapor pressure. H2O2 was produced by passing the carrier gases through a U-tube packed with H2O2−urea adduct (Sigma-Aldrich). Although the H2O2 concentration in the gas stream was not directly measured, [OH]0 was comparable to the HNO3 precursor method based on the intensity of the OH fluorescence signal. The rate constant measurements reported here were obtained using the HNO3 source, whereas the H2O2 source was used periodically to test for secondary reactions. Argon (99.999% purity), CO (Airgas, 1.03% mixed with Ar), and HNO3 or H2O2 were introduced into a 50 cm long mixing column attached to the chamber through calibrated MKS massflow controllers. Oxygen (Air Products UltraPure grade) was added in some experiments. The pressure in the chamber was measured with a 1000 Torr MKS Baratron capacitance manometer and stabilized using a MKS pressure control system. All flow controllers and Baratrons were calibrated against secondary standards that were accurate to better than 1% (1σ). Typical total flow rates were between 500 and 5000 sccm, resulting in an average residence time of ∼1 ms in the photolysis volume. This ensured that a fresh gas sample was photolyzed at each laser pulse. The reactor residence time was maintained constant over the pressure range 50−700 Torr. The temperature at the center of the reactor was monitored by a thermocouple and was stabilized within ±2 K. Gas concentrations were calculated using the measured flow rates, the total pressure, and the temperature in the cell. The O2 and CO concentration ranges were (0−5) × 1015 molecules cm−3 and (0.7−5) × 1015 molecules cm−3, respectively. The estimated uncertainty of the calculated [CO] was 2% (1σ), due mainly to the uncertainty in the CO mole fraction in the commercial gas cylinder. The initial OH concentration was approximately 4 orders of magnitude smaller ((3−6) × 1010 molecules cm−3) than [CO] to ensure first-order conditions and to minimize the contribution of fast secondary reactions. A copper shroud inside the chamber was used to control the gas temperature for the low-temperature studies. Cold nitrogen gas or liquid passed through copper tubing soldered to the shroud. The temperature of the gas mixture inside the chamber was controlled by heating the nitrogen coolant and was measured by a calibrated thermocouple located 3 mm downstream from the reaction region. Temperature fluctuates

EXPERIMENTAL METHOD

The OH + CO reaction rate constants were measured using the LP/LIF technique. A detailed description of the apparatus is given elsewhere.19 Only modifications for the current experiment and essential details are described here. Briefly, the photolysis and probe laser beams crossed at 90°. An 248 nm excimer laser (Lambda-Physik LPX120i, repetition rate of 20 Hz; ∼70 mW) was used to photolyze the OH precursors and generate OH radicals. A dye laser (Sirah; repetition rate of 20 kHz; ∼2 mW) pumped by a solid-state laser (Spectra-Physics YHP-40) was tuned to 282.0 nm to excite the P1 (1) line in the (1,0) band of the (A2Σ+, υ′ = 1) ← (X2Π, υ″ = 0) transition of the OH radical. Following rapid OH vibrational relaxation in the A state, fluorescence was monitored at 308 nm in the (0,0) band. The resulting OH fluorescence at 308 nm was collimated by a fast quartz lens, passed through a series of baffles and a high-transmission 308 nm interference filter, and collected by an Electron Tubes P25PC-02 photomultiplier tube (PMT), interfaced to a photon counting system. The OH was generated by photolysis of OH precursors, HNO320 or H2O2:21 HNO3 → OH + NO2 (248 nm, φ ≈ 0.88) B

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The Journal of Physical Chemistry A within ±2 K from the set point temperature. A computer equipped with a multichannel scaler card recorded the timedependent photon counts in 50 μs bins. OH decay signals were accumulated for 1000−5000 photolysis laser shots at each CO concentration.

for obtaining reliable rate constants was 193 K. Measurements at low temperatures were more difficult when using N2 due to its larger quenching of the OH fluorescence. As a result, measurements in N2 ranged from 223 to 296 K. The reaction rate constants were measured as a function of bath gas concentration at various temperatures as shown in Figure 2. At all temperatures, the rate constants increased with pressure in both N2 and Ar. The measured OH + CO rate constants are given in Table S1 with their statistical uncertainties and experimental conditions. The uncertainties range from 0.4 to 5.2%, 1σ, precision only. Additional contributions to the total error include the uncertainty in the mole fraction of the supplied CO mixture (2%), calibration of the pressure gauge (1%), and calibration of the flow controllers (2%). The estimated total uncertainties from random and systematic errors are ∼5−10%, 1σ. To study the possible effect of added O2 on the OH + CO reaction, O2 concentrations in the range (0−6) × 1015 molecules cm−3 were added to the reactant mixture at 296 K and 600−750 Torr of Ar. With added O2, [OH] vs time profiles could not be fitted with a single exponential function although satisfactory fits were obtained with a biexponential decay function as shown in Figure 3. Similar results were obtained when OH was produced using H2O2 photolysis. Figure 3 shows a comparison of least-squares fits to the OH data using exponential and biexponential functions. The residual curves (Figure 3b,c) show that the biexponential fit best describes the OH decay profile when O2 is present, whereas a single exponential is sufficient to fit the OH decay in the absence of added O2 (Figure 3d,e).

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RESULTS To determine the bimolecular rate constants, each OH temporal profile was fitted to an exponential function to obtain a pseudo-first-order rate constant, k′, from the following equation: OH signal = A · exp( − k′t ) + background

Typically, first-order OH decays were measured at 5−10 different CO concentrations for each temperature and pressure. The k′ values were plotted as a function of the corresponding CO concentration. Figure 1 shows the k′ values at a few

Figure 1. Fitted first-order decay constant, k′ vs [CO], at a total pressure of 200 Torr (■), 300 Torr (◆), and 700 Torr (▲) in N2 at 296 K.

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pressures at 296 K. Plots of k′ vs [CO] had y-intercepts in the range 90−900 s−1. This intercept is attributable to removal of OH radicals by reaction with the HNO3 precursor, and diffusion outside the viewing zone. The bimolecular reaction rate constant, k1, was determined from the slope of the linear least-squares fits through the ([CO], k′) data points. Rate constants for reaction 1 were measured over the temperature range 193−298 K for M = Ar and 223−298 K for M = N2. At low temperature, [OH]0 was limited due to condensation of the OH precursor, HNO3, on the cell walls. When using Ar as the carrier gas, the lowest useful temperature

DISCUSSION

A. Pressure and Temperature Dependences for Reaction 1. As discussed above, the OH + CO reaction involves the formation of a vibrationally excited HOCO intermediate, which decomposes to reactants (OH + CO) and products (H + CO2) or undergoes collisional stabilization followed by subsequent reaction with O2 (if present) to form HO2 + CO2. Equation 1 describes the overall rate constant as the sum of both bimolecular and termolecular terms using a modified Troe falloff function:2,11

Figure 2. Rate constants, k, for the OH + CO reaction as a function of [M] at various temperatures in (a) Ar and (b) N2. Error bars shown are precision only (1σ). (a) M = Ar: T = 193 K (▼), 223 K (●), 243 K (▲), 273 K (■), 296 K (◆). (b) M = N2: T = 223 K (■), 243 K (●), 265 K (▲), 288 K (▼), 296 K (◆). C

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Figure 3. (a) OH signal profile in Ar at 700 Torr with [O2] = 1.3 × 1014 molecules cm−3 and [CO] = 2.45 × 1015 molecules cm−3. The blue line is the biexponential fit, and the red line is the exponential fit. The fits include a constant term to account for the PMT background signal. (b) Data residue vs time plot from an exponential fit to data in panel (a). (c) Data residue vs time plot from biexponential fit to data in panel (a). (d) OH signal profile in Ar at 700 Torr without O2 at [CO] = 2.04 × 1015 molecules cm−3. (e) Data residue vs time plot from exponential fit to data in panel (d).

k = k bi(T ) +

k 0(T )[M] 1+

k 0(T )[M] k∞(T )

2 −1

× 0.6{1 + [log10(k 0(T )[M]/ k∞(T ))] }

(1)

where

k bi(T ) = A e−(Ea / RT ) ⎛ T ⎞−n ⎟ k 0(T ) = k 0300⎜ ⎝ 300 ⎠

(2) −m T ⎞ 300⎛ ⎜ ⎟ k∞(T ) = k∞ ⎝ 300 ⎠

(3)

In eq 2, A is the pre-exponential factor and Ea is the activation energy for the pressure independent bimolecular rate constant. In eq 3, k0(T) and k∞(T) are the low- and highpressure limiting rate constants at temperature T derived from measured values at 300 K. The parameters n and m are derived from temperature dependence data fits to falloff curves. The experimental pressure range in this study (50−700 Torr) is far below the pressure of the observed high-pressure limit, which is over 100 bar;16 therefore, it is difficult to derive an accurate value of k∞(T) from fits to experimental falloff curves. In this study, at a given temperature, kbi(T) was the intercept obtained from the fit and k0(T) was obtained from fits to falloff data, using values of k∞(T) and m from the NASA recommendations.1 Figure 4 shows an example of the fit to falloff data at 223 K. With k∞(223 K) and m fixed at the NASA recommended values of 7.48 × 10−13 cm3 molecule−1 s−1 and −1.3, respectively, the best-fit values of kbi (223 K) and k0 (223

Figure 4. Fall-off curve fitted to rate constant data for M = Ar at 223 K. . Error bars shown are precision only (1σ).

K) were (1.47 ± 0.03) × 10−13 cm3 molecule−1 s−1 and (6.1 ± 0.5) × 10−33 cm6 molecule−2 s−1, respectively. Using the fitted kbi, an Arrhenius plot of this bimolecular reaction term of eq 1 is obtained as shown in Figure 5. The derived values for the A-factor and Ea/R are (1.54 ± 0.14) × 10−13 cm3 molecule−1 s−1 and 13 ± 17 K, respectively, with an error of one standard deviation. Our calculated Ea/R produced a negligible positive temperature dependence for the bimolecular channel. However, this positive trend is not significant considering its error limits, as shown in Figure 5. The calculated Ea/R is 13 K and the standard deviation is 17 K, which is a factor of 1.3 larger than the Ea/R value. The very D

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Figure 5. Arrhenius plot of bimolecular reaction component of CO + OH reaction. Error limits are one standard deviation. Figure 6. Facsimile simulation of OH time profile with different concentration of added O2 at room temperature and 600 Torr. Black line: [O2] = 0 molecules cm−3, red line: [O2] = 1 × 1014 molecules cm−3, blue line: [O2] = 3 × 1014 molecules cm−3, Green line: [O2] = 5 × 1014 molecules cm−3, purple line [O2] = 1 × 1015 molecules cm−3, yellow line [O2] = 5 × 1015 molecules cm−3.

small temperature dependence is consistent with the results of Senosiain et al.,17 McCabe et al.,2 and the NASA1 and IUPAC18 recommendations. The fitted falloff parameters are listed in Table 1 for the termolecular component (second term in eq 1). The results show that k0300 for M = N2 is about 1.8 times larger than for M = Ar due to more efficient collisional energy transfer in N2 compared with Ar. The measured temperature dependence parameters for N2 and Ar are the same within the experimental uncertainties (n = 1.9 ± 0.5 and 2.0 ± 0.4, respectively). Both values are in reasonable agreement with the value for M = Air (n = 1.4) recommended in the NASA data evaluation.1 Table 1 gives a summary of current literature data for temperatures lower than 298 K. In this table, rate constants from this work and the NASA evaluation used eq 1 to parametrize the temperature dependence of the termolecular component, whereas in the IUPAC recommendation18 and McCabe et al.,2 the CO + OH rate constants were assumed to be temperature-independent. B. Effect of Added O2. In experiments with added [O2] greater than about 1014 molecules cm−3, the OH temporal profiles slightly departed from single-exponential decay. Under these conditions, a biexponential function provided a satisfactory fit to the OH data. This behavior suggests that with O2 present, there is a secondary source of OH. To gain insight into the mechanism, the Facsimile simulation program was used to integrate a kinetic mechanism consisting of 13 reactions. Table 2 lists the reactions in the mechanism and their

rate constants under our experimental condition. The initial concentration of OH is set as 3.5 × 1010 molecules cm−3, which was calculated from known precursor concentrations and the measured laser fluence. To simulate a range of typical experimental conditions, the initial concentrations of O2 were set to 0, 1 × 1014, 3 × 1014, 5 × 1014, 1 × 1015, and 5 × 1015 molecules cm−3. Results of the simulations are plotted in Figure 6. There are two possible sources of secondary OH. One source is the thermal decomposition of HOCO back to OH + CO (reaction 2b), in competition with the reaction of HOCO with O2 (reaction 6). However, on the basis of the measurements of HOCO decomposition rates by Fulle et al.,16 this process may be neglected at and below 298 K. The second, and mostly likely, OH source involves the secondary reactions of H produced in reaction 1: H + O2 + M → HO2 + M

(12)

H + HO2 → 2OH

(7)

In the absence of O2, HO2 and therefore secondary OH are not formed. In the limit of large [O2] there is insufficient H remaining to drive reaction 7. OH decays show single-

Table 2. Reactions and Rate Constants Used in Kinetics Simulationsa reaction

rate constant

ref.

OH + CO → CO2 + H (1) OH + CO → HOCO (2a) HOCO → OH + CO + M (2b) OH + OH → O + H2O (3) O + OH → O2 + H (4) OH + NO2 → product (5) HOCO + O2 → HO2 + CO2 (6) H + HO2 → OH + OH (7) H + HO2 → O + H2O (8) H + HO2 → H2 + O2 (9) O + HO2 → OH + O2 (10) OH + H2O2 → H2O + HO2 (11) H + O2 + M→ HO2 + M (12)

k1 = 1.54 × 10−13 k2a = 8 × 10−14 k2b = 3.7 × 10−11 k3 = 1.8 × 10−12 k4 = 2.2 × 10−11 × exp(120/T) k5 = 1.1 × 10−11 k6 = 2.0 × 10−12 k7 = 7.2 × 10−11 k8 = 1.6 × 10−12 k9 = 6.9 × 10−12 k10 = (3.0 × 10−11) × exp(200/T) k11 = 1.8 × 10−12 k12 = 7.4 × 10−13

this work this work 16 1 1 1 1 1 1 1 1 1 1

The term molecular reaction rates are calculated at 600 Torr when M = N2. Units for rate constants: first order, s−1; second order, cm3 molecule−1 s−1; third order, cm6 molecule−2 s−1. a

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exponential behavior under these conditions. With intermediate O2 concentrations, reaction 7 produces sufficient OH on the time scale of reaction 1 to result in biexponential OH decay. In Figure 6, the plots with [O2] = 1 × 1014, 3 × 1014, 5 × 1014, and 1 × 1015 molecules cm−3 show biexponential decay, whereas O2 concentrations larger than 5 × 1015 molecules cm−3 show no apparent biexponential decay. Figure 6 shows that although the OH temporal profiles differ significantly after about 6 decay lifetimes, the initial decay profiles are identical. Therefore, if we derive rate constants only from the initial decay profiles, the results should be independent of [O2]. For experiments with [O2] > 1 × 1014 molecules cm−3, the biexponential function

Article

ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpca.5b07220. Summary of all the experimental conditions and the rate constants (PDF)

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AUTHOR INFORMATION

Corresponding Author

*S. P. Sander. E-mail: [email protected]. Phone: (818) 354-2625. Fax: (818) 393-5065. Notes

The authors declare no competing financial interest.

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OH signal = A e−a1t + Be−a2t + C

ACKNOWLEDGMENTS This work was supported by the NASA Mars Fundamental Research, Upper Atmosphere Research and Tropospheric Chemistry Programs. The authors thank A. Komissarov who carried out initial experimental studies, members of the JPL Lab Studies and Modeling Group for helpful discussions, and Dave Natzic for technical assistance. Copyright 2015, California Institute of Technology.

was used to fit the OH temporal profiles, and the first-order rate constants for the OH + CO reaction were identified with the constant a1. With this approach, the overall rate constants are independent of [O2], as shown in Figure 7.

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REFERENCES

(1) Sander, S. P. A.; Barker, J.; Burkholder, J. R.; Friedl, J. B.; Golden, R. R.; Huie, D. M.; Kolb, R. E.; Kurylo, C. E.; Moortgat, M. J. Chemical Kinetics and Photochemical Data for Use in Atmospheric Studies, Evaluation Number 17; JPL Publication No. 10-6; Jet Propulsion Laboratory: Pasadena, CA, 2011 (2) McCabe, D. C.; Gierczak, T.; Talukdar, R. K.; Ravishankara, A. R. Kinetics of the Reaction OH plus CO under Atmospheric Conditions. Gophys. Res. lett. 2001, 28 (16), 3135−3138. (3) Krasnopolsky, V. A. Photochemistry of the Martian Atmosphere: Seasonal, Latitudinal, and Diurnal Variations. Icarus 2006, 185 (1), 153−170. (4) Atreya, S. K.; Gu, Z. G. Stability of the Martian Atmosphere: Is Heterogeneous Catalysis Essential? J. Geophys. Res. 1994, 99 (E6), 13133−13145. (5) Paraskevopoulos, G.; Irwin, R. S. Rates of OH Radical Reactions 0.11. The Pressure-Dependence of the Rate-Constant of the Reaction of OH Radicals with CO. J. Chem. Phys. 1984, 80 (1), 259−266. (6) Perry, R. A.; Atkinson, R.; Pitts, J. N. Kinetics of Reactions of OH Radicals with C2H2 and CO. J. Chem. Phys. 1977, 67 (12), 5577−5584. (7) Chan, W. H.; Uselman, W. M.; Calvert, J. G.; Shaw, J. H. Pressure-dependence of Rate Constant for Reaction OH+CO→ H +CO2. Chem. Phys. Lett. 1977, 45 (2), 240−244. (8) Biermann, H. W.; Zetzsch, C.; Stuhl, F. Pessure-Dependence of Reaction of HO with CO. Ber. Bunsen-Ges. Phys. Chem. 1978, 82 (6), 633−639. (9) Cox, R. A.; Derwent, R. G.; Holt, P. M. Relative Rate Constants for Reactions of OH Radicals with H2, CH4, CO, NO and HONO at Atmospheric-Pressure and 296 K. J. Chem. Soc., Faraday Trans. 1 1976, 72, 2031−2043. (10) Butler, R.; Solomon, I. J.; Snelson, A. Pressure-Dependence of CO + OH Rate Constant in O2 + N2 Mixtures. Chem. Phys. Lett. 1978, 54 (1), 19−24. (11) Demore, W. B. Rate-constant for the OH + CO Reaction Pressure Dependence and the Effect of Oxygen. Int. J. Chem. Kinet. 1984, 16 (10), 1187−1200. (12) Hofzumahaus, A.; Stuhl, F. Rate-Constant of the Reaction HO + CO in the Presence of N2 and O2. Ber. Bunsen-Ges. Phys. Chem. 1984, 88 (6), 557−561. (13) Hynes, A. J.; Wine, P. H.; Ravishankara, A. R. Kinetics of the OH + CO Reaction under Atmospheric Conditions. J. Geophys. Res. 1986, 91 (D11), 11815−11820.

Figure 7. CO + OH rate constants for experiment with and without O2 at room temperature (296 K). The red circle is the data without O2, and the black square is the data with [O2] = 5 × 1015 molecules cm−3; the error bar is one standard deviation, σk.

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CONCLUSIONS In this study, the kinetics of the OH + CO reaction are studied over the temperature range 193−296 K and pressure range 50− 750 Torr for M = N2 and Ar. Rate constants measured in our study are consistent with the NASA/JPL Data Evaluation.1 The rate constants can be expressed by a combination of a bimolecular and termolecular rate constants. For the bimolecular component, no temperature dependence is observed. For the termolecular component, the rate constant and temperature dependence in the low-pressure limit, k0(T), were determined using constraints on the high-pressure limit, k∞(T), from the NASA Data Evaluation.1 The low-pressure limit rate constant in N2 is about a factor of 1.8 larger than in Ar. Laboratory studies and model simulations of the effect of added O2 are carried out. These results show that O2 does not directly affect the CO + OH reaction rate constant measurements. This study reduces the uncertainties in previously evaluated rate constants for the OH + CO reaction under conditions representative of Earth’s troposphere and stratosphere, and the atmosphere of Mars. F

DOI: 10.1021/acs.jpca.5b07220 J. Phys. Chem. A XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry A (14) Miyoshi, A.; Matsui, H.; Washida, N. Detection and Reactions of the HOCO Radical in Gas-Phase. J. Chem. Phys. 1994, 100 (5), 3532−3539. (15) Frost, M. J.; Sharkey, P.; Smith, I. W. M. Reaction between Hydroxyl (Deuteroxyl) Radicals and Carbon Monoxide at Temperatures down to 80 K: Experiment and Theory. J. Phys. Chem. 1993, 97 (47), 12254−12259. (16) Fulle, D.; Hamann, H. F.; Hippler, H.; Troe, J. High Pressure Range of Addition Reactions of HO. II. Temperature and Pressure Dependence of the Reaction HO+CO → HOCO → H+CO2. J. Chem. Phys. 1996, 105 (3), 983−1000. (17) Senosiain, J. P.; Musgrave, C. B.; Golden, D. M. Temperature and Pressure Dependence of the Reaction of OH and CO: Master Equation Modeling on a High-level Potential Energy Surface. Int. J. Chem. Kinet. 2003, 35 (9), 464−474. (18) Atkinson, R.; Baulch, D. L.; Cox, R. A.; Crowley, J. N.; Hampson, R. F.; Hynes, R. G.; Jenkin, M. E.; Rossi, M. J.; Troe, J.; IUPAC Subcommittee. Evaluated kinetic and Photochemical Data for Atmospheric Chemistry: Volume II - Gas Phase Reactions of Organic Species. Atmos. Chem. Phys. 2006, 6 (11), 3625−4055. (19) Mollner, A. K.; Valluvadasan, S.; Feng, L.; Sprague, M. K.; Okumura, M.; Milligan, D. B.; Bloss, W. J.; Sander, S. P.; Martien, P. T.; Harley, R. A.; Mccoy, A. B.; Carter, W. P. L. Rate of Gas Phase Association of Hydroxyl Radical and Nitrogen Dioxide. Science 2010, 330, 646−649. (20) Riffault, V.; Gierczak, T.; Burkholder, J. B.; Ravishankara, A. R. Quantum Yields for OH Production in the Photodissociation of HNO3 at 248 and 308 nm and H2O2 at 308 and 320 nm. Phys. Chem. Chem. Phys. 2006, 8 (9), 1079−1085. (21) Volman, D. H. The Vapor-Phase Photo Decomposition of Hydrogen Peroxide. J. Chem. Phys. 1949, 17 (10), 947−950.

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DOI: 10.1021/acs.jpca.5b07220 J. Phys. Chem. A XXXX, XXX, XXX−XXX