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Hydration Leads to Efficient Reactions of the Carbonate Radical Anion with Hydrogen Chloride in the Gas Phase Wai Kit Tang,† Christian van der Linde,‡ Chi-Kit Siu,*,† and Martin K. Beyer*,‡ †

Department of Biology and Chemistry, City University of Hong Kong, 83 Tat Chee Avenue, Kowloon Tong, Hong Kong SAR, China ‡ Institut für Ionenphysik und Angewandte Physik, Leopold-Franzens-Universität Innsbruck, Technikerstraße 25, 6020 Innsbruck, Austria S Supporting Information *

ABSTRACT: The carbonate radical anion CO3•− is a key intermediate in tropospheric anion chemistry. Despite its radical character, only a small number of reactions have been reported in the literature. Here we investigate the gas-phase reactions of CO3•− and CO3•−(H2O) with HCl under ultrahigh vacuum conditions. Bare CO3•− forms OHCl•− with a rate constant of 4.2 × 10−12 cm3 s−1, which corresponds to an efficiency of only 0.4%. Hydration accelerates the reaction, and ligand exchange of H2O against HCl proceeds with a rate of 2.7 × 10−10 cm3 s−1. Quantum chemical calculations reveal that OHCl•− is best described as an OH• hydrogen bonded to Cl−, while the ligand exchange product is Cl−(HCO3•). Under tropospheric conditions, where CO3•−(H2O) is the dominant species, Cl−(HCO3•) is efficiently formed. These reactions must be included in models of tropospheric anion chemistry.

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et al.,23 who established an upper limit of 3 × 10−11 cm3 s−1 for the rate constant. Consequently, reactions of CO3•− with HCl are not considered in models of tropospheric anion chemistry.1,26,28 Below 3 km altitude, HCl mixing ratios of 50−500 pptv have been reported,29 comparable to HNO3.30 Given this known reactivity of various organic and inorganic acids, the relatively high upper limit for the reactivity of HCl, and the importance of HCl as a tropospheric trace gas, we found it worthwhile to have a closer look at this reaction. We therefore investigated the gas-phase reaction of CO3•− and CO3•−(H2O) with HCl by a combination of Fourier transform ion cyclotron resonance (FTICR) mass spectrometry27,31,32 and quantum chemical calculations at the M06-2X/6-311++G(d,p) level of theory. Due to the high mass resolution and long trapping times, FTICR is ideally suited to study slow reactions of radical anions in the gas phase.

INTRODUCTION The carbonate radical anion CO3•− is a key intermediate in tropospheric anion chemistry.1 Fehsenfeld, Ferguson, and coworkers have shown2 that it is formed in the presence of electrons and ozone by electron attachment of an electron to O2, reaction 1, followed by charge transfer to ozone, reaction 2. The resulting ozonide transfers an O•− ion to CO2, yielding the CO3•− radical anion, reaction 3. Spectroscopically, CO3•− was characterized by matrix isolation infrared spectroscopy3 and photodissociation.4−7 Johnson, Viggiano, and co-workers established a O2C−O•− bond dissociation energy of 269 ± 5 kJ mol−1,8−10 by photodissociation and high-level quantum chemical calculations. CO3•−(H2O) was studied by photodissociation11−16 and high-pressure mass spectrometry.17 e− + O2 + M → O2•− + M,

M = O2 , N2 , Ar, H 2O, ... (1)

O2

•−

O3•−

+ O3 → O2 +

+ CO2 →

O3•−

CO3•−

+ O2

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EXPERIMENTAL AND COMPUTATIONAL DETAILS The experiments were performed on a modified Bruker/ Spectrospin CMS47X FT-ICR mass spectrometer described in detail before.31,32 CO3•− and CO3•−(H2O) ions are formed in a laser vaporization source.33−35 Metallic zinc is vaporized with a frequency-doubled Nd:YAG laser. The hot plasma is expanded in helium seeded with CO2 and O2. Due to the supersonic

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Despite its reactive radical character, only a small number of gas-phase reactions of CO3•− has been reported,18−22 in particular the reactions with nitrogen oxides that lead to the formation of nitrate NO3−.23−26 In a recent study of CO3•− with HCOOH,27 we observed a quite efficient reaction with a rate constant of 3.6 ± 0.9 × 10−10 cm3 s−1, which starts with a proton transfer from the acid molecule to CO3•−. Also nitric acid20−22 and methanesulfonic acid19 react with CO3•−, while no products could be identified in reactions with HCl by Dotan © XXXX American Chemical Society

Received: September 26, 2016 Revised: December 13, 2016 Published: December 13, 2016 A

DOI: 10.1021/acs.jpca.6b09715 J. Phys. Chem. A XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry A

above the noise level. Stoichiometrically, HCl2− and Cl2•− cannot be formed in a binary collision of CO3•− with HCl. The presence of (HCl)2 in the ICR cell, however, is highly improbable. The most realistic interpretation of this finding is that HCl2− and Cl2•− are secondary products of an intermediate that is formed very slowly. Based on the observation of traces of OHCl•−, we formulate reactions 4−6 as a plausible scenario.

expansion into high vacuum, the ions are rotationally and vibrationally cooled below room temperature, and electronically excited states can be ruled out. The ions are transferred to the ICR cell by a system of electrostatic lenses through several differential pumping states. In the ICR cell, CO3•− or CO3•−(H2O) is mass selected by resonant ejection of unwanted ions. HCl is introduced at constant backing pressure in the range of 10−8 mbar through a leak valve. Mass spectra are recorded after different reaction delays to monitor the reaction kinetics. Pressure is calibrated as described before.36,37 To extract kinetic parameters, the data are fitted assuming a pseudo-first order rate law with a genetic algorithm. Efficiencies Φ are calculated as the ratio of the pressure-corrected rate kabs and the collision rate kADO estimated from average dipole orientation (ADO) theory,38 Φ = kabs/kADO. The pressure calibration is the dominant source of error and is expected to be around 25%. The noise level in the kinetics plots is the noise level of the mass spectra taken at the respective reaction delay, normalized to the total intensity of the ions in the fit. Theoretical reaction profiles were obtained using quantum chemical calculations at the M06-2X/6-311++G(d,p) level of theory, employing the Gaussian09 program package.39 This theory level produced results in agreement with literature thermochemistry in our recent study on the reactions of CO3•− with HCOOH.27 The energies of all optimized geometries were corrected with zero-point energies calculated from harmonic vibration analyses.

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CO3•− + HCl → OHCl•− + CO2

(4)

OHCl•− + HCl → HCl 2− + OH•

(5)

OHCl•− + HCl → Cl 2•− + H 2O

(6)

To extract the kinetics of the reaction, we recorded mass spectra at reaction delays from 0 to 30 s. Figure 2a shows a fit

RESULTS AND DISCUSSION

Figure 1 shows mass spectra of the reaction of mass selected CO3•− at 0, 5, and 30 s reaction delay. The dominant product species are HCl2− and Cl2•−, with traces of OHCl•− hardly

Figure 2. Kinetics of the reaction of (a) CO3•− with HCl, employing reactions 4−6, and (b) CO3•−(H2O) with HCl, employing reactions 4−9. Normalized intensities from the mass spectra are shown as circles. The lines denote a fit obtained with a genetic algorithm.

of the kinetics with reactions 4−6. The fit works extremely well, corroborating that reactions 4−6 take place. Pressure-dependent pseudo-first order rate constants from the fit have been converted to pressure-independent bimolecular rate constants and efficiencies,36−38 which are summarized in Table 1. In Table 1. Room Temperature Rate Constants kabs and Efficiencies Φ = kabs/kADO of Reactions 4−6a kabs (cm3 s−1)

reaction •−

•−

CO3 + HCl → OHCl + CO2 (4) OHCl•− + HCl → HCl2− + OH• (5) OHCl•− + HCl → Cl2•− + H2O (6) CO3•−(H2O) + HCl → [CO3,HCl]•− + H2O (7) [CO3,HCl]•− + HCl → [CO3,HCl,H]− + Cl• (9)

Figure 1. Mass spectra of the reaction of CO3•− with HCl at a pressure of 2.4 × 10−8 mbar. (a) Mass selected CO3•− at 0 s reaction delay. Main product species HCl2− and Cl2•− and traces of OHCl•− after (b) 5 s and (c) 30 s.

a

B

4.2 1.5 1.5 2.7 1.5

× × × × ×

−12

10 10−9 10−10 10−10 10−11

Φ (%) 0.4 122 12 24 1.4

The uncertainty of the values is ±25%. DOI: 10.1021/acs.jpca.6b09715 J. Phys. Chem. A XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry A agreement with the work by Dotan et al., the rate constant of reaction 4 with 4.2 × 10−12 cm3 s−1 lies below the previously established upper limit, only one out of 250 collisions is reactive. The combined efficiency of 134% of reactions 5 and 6 indicates that in the second step, every collision between the anionic OHCl•− and the highly polar HCl leads to a reaction, if the HCl pressure is underestimated by 25%. This corresponds to the uncertainty of the pressure measurement. One might expect that the kinetic parameters for reactions 5 and 6 are very sensitive to the slow initial reaction and the low concentration of OHCl•− and therefore might have fairly large uncertainties. This would indeed be the case if the signal of the intermediate OHCl•− would be associated with a large error. However, we worked hard to improve the total signal and to adjust the HCl pressure so that the intermediate is observed with reliable intensity throughout the kinetics. The noise on the data points in Figure 2 is remarkably low. Any change in the parameters of reactions 5 and 6 would translate into a proportional parallel shift of the blue line in Figure 2, leading to significant deviations from the data. Also the rate of the slow first step is reliable: it is directly given by the decay of the CO3•− intensity, which is reproduced perfectly by the fit. Even under these overall unfavorable conditions, we can conclude that all rates are reliable, subject to the 25% error margin mentioned above. In the troposphere, however, CO3•−(H2O) is the dominant species up to an altitude of 11 km under typical temperature and humidity conditions.26 Near the ground, CO3•−(H2O) constitutes 0.9% of the total anion concentration.26 In our mass spectrometer, CO3•−(H2O) undergoes ligand exchange with HCl, reaction 7, accompanied by blackbody infrared radiative dissociation of CO3•−(H2O), reaction 8. The ligand exchange product [CO3, HCl]•− abstracts a hydrogen atom from HCl in a secondary reaction 9. Figure 2b shows the kinetics of the reaction; mass spectra are available as Supporting Information. The rate of the ligand exchange reaction 7 is 2.7 × 10−10 cm3 s−1, i.e., hydration increases the reactivity of CO3•− with HCl by almost 2 orders of magnitude. CO3•−(H 2O) + HCl → [CO3 , HCl]•− + H 2O

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CO3•−(H 2O) → CO3•− + H 2O

(8)

[CO3 , HCl]•− + HCl → [CO3 , HCl, H]− + Cl•

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Figure 3. Potential energy surfaces of (a) reactions 4 (black) and 7 and 8 (green); (b) reactions 5 (blue) and 6 (red). Energies with zeropoint correction, in kJ mol−1, were evaluated at the M06-2X/6-311+ +G(d,p) level of theory. Some selected bond distances, in Å, are shown. Spin density is shown with an isovalue of 0.01 au.

Interestingly, we found another local minimum 1′ in the entrance channel, in which HCl coordinates to an oxygen atom of CO3•− with its chlorine atom. Without zero-point correction, 1′ lies 2 kJ mol−1 below the separated reactants. Zero-point effects shift this local minimum up to 4 kJ mol−1 above the entrance channel. This stationary point on the potential energy surface is connected to complex 2 via TS1′, which again due to zero-point effects is isoenergetic with 1′. The consequence of the local minimum 1′ is that trajectories exist on which HCl molecules approaching CO3•− with the chlorine end experience a repulsive interaction. This qualitatively rationalizes why the reaction proceeds with a rate below the collision rate. A quantitative analysis would require multidimensional classical trajectory calculations of the CO3•− + HCl collision,40,41 which goes far beyond the scope of the present study. The second reaction step, Figure 3b, starts with formation of a ternary hydrogen bonded complex 6 between OH•, Cl−, and HCl, at −83 kJ mol−1 relative to the OHCl•− + HCl entrance channel 5. From there, loss of OH• completes reaction 5, which is basically a ligand exchange of OH• against HCl, with an overall energy of the products 7 of −40 kJ mol−1. The HCl2− is a hydrogen bonded complex of Cl−···H+···Cl−. Interestingly, the proton of HCl in 6 can transfer to OH• via the transition structure TS2 at −38 kJ mol−1, which is almost isoenergetic to the release of OH• also from 6. H2O formation is highly exothermic, resulting in a Cl−···H2O···Cl• intermediate 8, in which a hydrogen atom of H2O is hydrogen bonded to Cl− with the H···Cl distance of 1.91 Å, and the oxygen atom of H2O

To learn more about the reaction mechanism and the structure of the products, we calculated the potential energy surface of reactions 4−6 at the M06-2X/6-311++G(d,p) level of theory. Figure 3a shows the results for the first step, reaction 4. The proton is transferred already in the initial complex 2, which can be described as CO3H• radical hydrogen bonded to Cl− with the H···Cl distance of 1.75 Å, lying at −109 kJ mol−1, far below the entrance channel 1. Passing the barrier to C−O bond breaking via TS1 at −34 kJ mol−1 results in a ternary complex 3 of CO2, OH•, and Cl− at an energy of −61 kJ mol−1. In TS1, the spin has significantly localized already on the incipient OH•. The proton assists largely the C−O breaking with a barrier of 75 kJ mol−1, significantly lower than that for the bare CO3•− (269 ± 5 kJ mol−1, vide supra). The binding energy of CO2 in 3 is only 23 kJ mol−1; therefore, loss of CO2 and formation of the OHCl•− intermediate 4 is significantly exothermic with respect to the separated reactants 1 by 38 kJ mol−1. The intermediate 4 can be described as an OH• radical hydrogen bonded to Cl− with the H···Cl distance of 2.05 Å. C

DOI: 10.1021/acs.jpca.6b09715 J. Phys. Chem. A XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry A forms a two-center-three-electron (2c3e) bond with Cl• with the Cl···O distance of 2.34 Å. With a moderate relative barrier of 22 kJ mol−1, the 2c3e bond between Cl• and OH2 in 8 can switch to Cl2•− via TS3, in which the Cl···Cl distance is shortened to 3.57 Å, and some spin density on the ClOH2• moiety is delocalized into Cl−. This rearrangement of the ternary complex 8 results in the significantly lower-energy binary complex 9, Cl2•−···H2O, in which the newly formed 2c3e radical-anion pair, Cl2•−, with the Cl···Cl distance of 2.65 Å is stabilized by H2O through two hydrogen bonds with the two H···Cl distances of 2.50 Å. Loss of water completes reaction 6, which is overall exothermic by 124 kJ mol−1. The calculations explain the occurrence of reactions 4−6. All transition states lie well below the entrance channels, and the secondary reactions are strongly exothermic. The branching ratio of 10:1 between reactions 5 and 6 is consistent with the tight transition state of H2O formation, reaction 6, in combination with the slightly higher activation energy, while loss of an OH• radical, reaction 5, involves a loose transition state. The binding energy of water in CO3•−(H2O), reaction 8, is calculated to be 62 kJ mol−1, very close to the value of 59 kJ mol−1 reported by Castleman and co-workers from highpressure mass spectrometry.17 Together with Figure 3a, this immediately explains reaction 7: formation of 2 is exothermic by 47 kJ mol−1, and the [CO3,HCl]•− product is best described as Cl−(HCO3•). Evaporation of H2O prevents the onward reaction via TS1 to 3. The hydrogen atom transfer (reaction 9) is calculated to be slightly exothermic by 5 kJ mol−1, explaining the small efficiency of this reaction. On the basis of reaction 5, one may speculate that in the troposphere, where HCl concentrations are low, the OHCl•− intermediate preferentially undergoes ligand exchange with abundant molecules, e.g., H2O. We therefore examine theoretically the thermochemistry of reaction 10 to check whether this scenario is plausible, Figure 4. OHCl•−(H 2O)n + H 2O → Cl−(H 2O)n + 1 + OH•

Given the large abundance of H2O in the troposphere, it is reasonable to expect that the multiple hydrates OHCl•−(H2O)n are formed in significant amounts. Successive hydrations of the OHCl•− core for n = 1−4 are exothermic by 51−58 kJ mol−1. Displacing the OH•, i.e., the ligand exchange reaction 10, is predicted to be only slightly endothermic by 15 kJ mol−1 for n = 0, which gradually decreases with hydration to 14 kJ mol−1 for n = 1 and 1 kJ mol−1 for n = 2, and then to 0 kJ mol−1 for n = 3. However, even for n = 0, the equilibrium will likely lie on the product side of reaction 10, given the high concentration of H2O and the low concentration of OH•. Therefore, it must be expected that OHCl•− formed in reaction 4 preferentially reacts onward. Reactions 1−4 followed by reaction 10 describe a reaction pathway in which HCl and electrons convert ozone to hydroxyl radicals. In the lower troposphere, the Cl−(HCO3•) radical 2 formed from CO3•−(H2O) is expected to react efficiently with saturated hydrocarbons. Combining our calculations with literature thermochemistry, H atom abstraction by 2 from ethane is 15 kJ mol−1 exothermic, and even H atom abstraction from methane is within reach, being only 2 kJ mol−1 endothermic.

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CONCLUSIONS Under the low pressure conditions in an FT-ICR mass spectrometer, we were able to identify the reaction pathway of CO3•− with HCl, which so far was thought to be unreactive. The reaction proceeds with a non-negligible efficiency of 0.4% to form OHCl•−, which undergoes fast secondary reactions to form Cl2•− and HCl2−. Interestingly, hydration accelerates the primary reaction by a factor of 60. Quantum chemical calculations show that, in the ligand exchange product [CO3•−, HCl], proton transfer is barrierless, resulting in the HCO3• radical hydrogen bonded to Cl−. The latter complex is an excellent hydrogen atom acceptor. The newly identified reactions of CO3•− and CO3•−(H2O) with HCl suggest that our current knowledge of tropospheric anion chemistry is only scratching the surface. In the lower troposphere, where CO3•−(H2O) dominates, reactions with HCl must be included in models of tropospheric anion chemistry. The reaction products are reactive radicals, which readily attack hydrocarbons.

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.jpca.6b09715. Mass spectra of the reaction of CO3•−(H2O); Cartesian coordinates, harmonic vibration analyses and energies of optimized structures shown in Figures 3 and 4 (PDF)

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AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected]. *E-mail: [email protected]. ORCID

Figure 4. Potential energy surface of reaction 10. Energies with zeropoint correction, in kJ mol−1, were evaluated at the M06-2X/6-311+ +G(d,p) level of theory. Ligand exchange energies (dashed lines) of OHCl•−(H2O)n + H2O → Cl−(H2O)n+1 + OH• are 15 kJ/mol for n = 0, 14 kJ/mol for n = 1, 1 kJ/mol for n = 2, and 0 kJ/mol for n = 3.

Martin K. Beyer: 0000-0001-9373-9266 Notes

The authors declare no competing financial interest. D

DOI: 10.1021/acs.jpca.6b09715 J. Phys. Chem. A XXXX, XXX, XXX−XXX

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(18) Dotan, I.; Davidson, J. A.; Streit, G. E.; Albritton, D. L.; Fehsenfeld, F. C. A Study of the Reaction O3−+CO2 ⇌ CO3−+O2 and its Implication on the Thermochemistry of CO3 and O3 and Their Negative Ions. J. Chem. Phys. 1977, 67, 2874−2879. (19) Schoon, N.; Amelynck, C.; Bultinck, P.; Arijs, E. Study of Ion/ Molecule Reactions of Atmospherically Important Negative Ions with Methane Sulfonic Acid. Int. J. Mass Spectrom. 2002, 221, 209−218. (20) Fehsenfeld, F. C.; Howard, C. J.; Schmeltekopf, A. L. Gas Phase Ion Chemistry of HNO3. J. Chem. Phys. 1975, 63, 2835−2841. (21) Möhler, O.; Arnold, F. Flow Reactor and Triple Quadrupole Mass Spectrometer Investigations of Negative Ion Reactions Involving Nitric Acid: Implications for Atmospheric HNO3 Detection by Chemical Ionization Mass Spectrometry. J. Atmos. Chem. 1991, 13, 33−61. (22) Shuman, N. S.; Hunton, D. E.; Viggiano, A. A. Ambient and Modified Atmospheric Ion Chemistry: From Top to Bottom. Chem. Rev. 2015, 115, 4542−4570. (23) Dotan, I.; Albritton, D. L.; Fehsenfeld, F. C.; Streit, G. E.; Ferguson, E. E. Rate Constants for the Reactions of O−, O2−, NO2−, CO3−, and CO4− with HCl and ClO− with NO, NO2, SO2, and CO2 at 300K. J. Chem. Phys. 1978, 68, 5414−5416. (24) Huertas, M.; Fontan, J. Formation of Stable Positive and Negative Small Ions of Tropospheric Interest. Atmos. Environ. 1982, 16, 2521−2527. (25) Kawamoto, H.; Ogawa, T. A Steady State Model of Negative Ions in the Lower Stratosphere. Planet. Space Sci. 1984, 32, 1223− 1233. (26) Kawamoto, H.; Ogawa, T. First Model of Negative Ion Composition in the Troposphere. Planet. Space Sci. 1986, 34, 1229− 1239. (27) van der Linde, C.; Tang, W. K.; Siu, C.-K.; Beyer, M. K. Electrons Mediate the Gas-Phase Oxidation of Formic Acid with Ozone. Chem. - Eur. J. 2016, 22, 12684−12687. (28) Luts, A.; Parts, T. Evolution of Negative Small Air Ions at Two Different Temperatures. J. Atmos. Sol.-Terr. Phys. 2002, 64, 763−774. (29) Vierkorn-Rudolph, B.; Bächmann, K.; Schwarz, B.; Meixner, F. X. Vertical Profiles of Hydrogen Chloride in the Troposphere. J. Atmos. Chem. 1984, 2, 47−63. (30) Cooper, M.; Martin, R. V.; Wespes, C.; Coheur, P.-F.; Clerbaux, C.; Murray, L. T. Tropospheric Nitric Acid Columns from the IASI Satellite Instrument Interpreted with a Chemical Transport Model: Implications for Parameterizations of Nitric Oxide Production by Lightning. J. Geophys. Res. Atmos 2014, 119, 10068−10079. (31) Berg, C.; Schindler, T.; Niedner-Schatteburg, G.; Bondybey, V. E. Reactions of Simple Hydrocarbons with Nbn+: Chemisorption and Physisorption on Ionized Niobium Clusters. J. Chem. Phys. 1995, 102, 4870−4884. (32) Höckendorf, R. F.; Balaj, O. P.; van der Linde, C.; Beyer, M. K. Thermochemistry from Ion−Molecule Reactions of Hydrated Ions in the Gas Phase: A New Variant of Nanocalorimetry Reveals Product Energy Partitioning. Phys. Chem. Chem. Phys. 2010, 12, 3772−3779. (33) Bondybey, V. E.; English, J. H. Laser Induced Fluorescence of Metal Clusters Produced by Laser Vaporization: Gas Phase Spectrum of Pb2. J. Chem. Phys. 1981, 74, 6978−6979. (34) Dietz, T. G.; Duncan, M. A.; Powers, D. E.; Smalley, R. E. Laser Production of Supersonic Metal Cluster Beams. J. Chem. Phys. 1981, 74, 6511−6512. (35) Maruyama, S.; Anderson, L. R.; Smalley, R. E. Direct Injection Supersonic Cluster Beam Source for FT-ICR Studies of Clusters. Rev. Sci. Instrum. 1990, 61, 3686−3693. (36) Bartmess, J. E.; Georgiadis, R. M. Empirical Methods for Determination of Ionization Gauge Relative Sensitivies for Different Gases. Vacuum 1983, 33, 149−153. (37) Akhgarnusch, A.; Hö c kendorf, R. F.; Beyer, M. K. Thermochemistry of the Reaction of SF6 with Gas-Phase Hydrated Electrons: A Benchmark for Nanocalorimetry. J. Phys. Chem. A 2015, 119, 9978−9985.

ACKNOWLEDGMENTS M.K.B. acknowledges startup funds from the University of Innsbruck. C.K.S. thanks City University of Hong Kong (CityU) for financial support (Project No. 7004401). W.K.T. acknowledges Chow Yei Ching School of Graduate Studies, CityU, for his postgraduate studentship and scholarship.

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REFERENCES

(1) Beig, G.; Brasseur, G. P. Model of Tropospheric Ion Composition: A First Attempt. J. Geophys. Res.-Atmos. 2000, 105, 22671−22684. (2) Adams, N. G.; Bohme, D. K.; Dunkin, D. B.; Fehsenfeld, F. C.; Ferguson, E. E. Flowing Afterglow Studies of Formation and Reactions of Cluster Ions of O2+, O2−, and O−. J. Chem. Phys. 1970, 52, 3133− 3140. (3) Jacox, M. E.; Milligan, D. E. Matrix-Isolation Study of the Vibrational Spectrum and Structure of the CO3− Radical Anion. J. Mol. Spectrosc. 1974, 52, 363−379. (4) Moseley, J. T. Photodissociation Spectroscopy of CO3−. J. Chem. Phys. 1976, 65, 2512−2517. (5) Beyer, R. A.; Vanderhoff, J. A. Cross Section Measurements for Photodetachment or Photodissociation of Ions Produced in Gaseous Mixtures of O2, CO2, and H2O. J. Chem. Phys. 1976, 65, 2313−2321. (6) Hunton, D. E.; Hofmann, M.; Lindeman, T. G.; Castleman, A. W. Photodissociation Dynamics of CO3−. J. Chem. Phys. 1985, 82, 134− 150. (7) Snodgrass, J. T.; Roehl, C. M.; van Koppen, P. A. M.; Palke, W. E.; Bowers, M. T. Photodissociation of CO3−: Product Kinetic Energy Measurements as a Probe of Excited State Potential Surfaces and Dissociation Dynamics. J. Chem. Phys. 1990, 92, 5935−5943. (8) Bopp, J. C.; Diken, E. G.; Headrick, J. M.; Roscioli, J. R.; Johnson, M. A.; Midey, A. J.; Viggiano, A. A. Determination of the CO3− Bond Strength via the Resonant Two-Photon Photodissociation Threshold: Electronic and Vibrational Spectroscopy of CO3−Arn. J. Chem. Phys. 2006, 124, 174302. (9) Viggiano, A. A. Reexamination of Ionospheric Chemistry: High Temperature Kinetics, Internal Energy Dependences, Unusual Isomers, and Corrections. Phys. Chem. Chem. Phys. 2006, 8, 2557− 2571. (10) Viggiano, A. A.; Midey, A. J.; Ehlerding, A. Kinetics of the Reactions of ONOO− with Small Molecules. Int. J. Mass Spectrom. 2006, 255, 65−70. (11) Smith, G. P.; Lee, L. C.; Cosby, P. C.; Peterson, J. R.; Moseley, J. T. Photodissociation and Photodetachment of Molecular Negative Ions. V. Atmospheric Ions from 7000 to 8400 Å. J. Chem. Phys. 1978, 68, 3818−3822. (12) Smith, G. P.; Lee, L. C.; Moseley, J. T. Photodissociation and Photodetachment of Molecular Negative Ions. VII. Ions Formed in CO2/O2/H2O mixtures, 3500−5300 Å. J. Chem. Phys. 1979, 71, 4034−4041. (13) Hunton, D. E.; Hofmann, M.; Lindeman, T. G.; Albertoni, C. R.; Castleman, A. W. Photodissociation Spectroscopy and Dynamics of Negative Ion Clusters. II. CO3−·(H2O)1,2,3. J. Chem. Phys. 1985, 82, 2884−2895. (14) Castleman, A. W.; Albertoni, C. R.; Marti, K.; Hunton, D. E.; Keesee, R. G. Photodissociation of Negative Ions and Their Clusters. Faraday Discuss. Chem. Soc. 1986, 82, 261−273. (15) Snodgrass, J. T.; Kim, H.-S.; Bowers, M. T. Intramolecular Energy Transfer Rates in Photoexcited Cluster Ions: The Photodissociation Dynamics of CO3−·H2O and CO3−·CO2. J. Chem. Phys. 1988, 88, 3072−3080. (16) Roehl, C. M.; Snodgrass, J. T.; Deakyne, C. A.; Bowers, M. T. Photodissociation of CO3−·H2O: Observation of the O−·H2O+CO2 Product Channel. J. Chem. Phys. 1991, 94, 6546−6552. (17) Keesee, R. G.; Lee, N.; Castleman, A. W. Properties of Clusters in the Gas Phase. 3. Hydration Complexes of CO3− and HCO3−. J. Am. Chem. Soc. 1979, 101, 2599−2604. E

DOI: 10.1021/acs.jpca.6b09715 J. Phys. Chem. A XXXX, XXX, XXX−XXX

Article

The Journal of Physical Chemistry A (38) Bass, L.; Su, T.; Chesnavich, W. J.; Bowers, M. T. Ion-Polar Molecule Collisions. A Modification of Average Dipole Orientation Theory: The cos θ Model. Chem. Phys. Lett. 1975, 34, 119−122. (39) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Scalmani, G.; Barone, V.; Mennucci, B.; Petersson, G. A.; et al. Gaussian 09, revision D.01; Gaussian, Inc.: Wallingford, CT, 2013. (40) Lourderaj, U.; Hase, W. L. Theoretical and Computational Studies of Non-RRKM Unimolecular Dynamics. J. Phys. Chem. A 2009, 113, 2236−2253. (41) Xie, J.; Otto, R.; Mikosch, J.; Zhang, J.; Wester, R.; Hase, W. L. Identification of Atomic-Level Mechanisms for Gas-Phase X− + CH3Y SN2 Reactions by Combined Experiments and Simulations. Acc. Chem. Res. 2014, 47, 2960−2969.

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