Measurement of the rate constants for the reactions of the IO• radical with sulfur-containing compounds H2S, (CH3)2S, and SO2

Larin, I. K.; Messineva, N. A.; Spasskii, A. I.; Trofimova, E. M.; Turkin, L. E.

doi:10.1007/bf02756058

Cited by 15 publications

(8 citation statements)

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“…Although the reaction between IO and SO 2 has a small barrier (7.3 kJ mol −1 ), the reaction has quite a tight transition state and the TST calculation yields k(200-400 K) = 4.3 × 10 −14 exp(−1150 / T ) cm 3 s −1 . The resulting value of k(343 K) = 1.6 × 10 −15 cm 3 s −1 is consistent with an experimental upper limit of 5.6 × 10 −15 cm 3 s −1 determined at that temperature by Larin et al (2000). At a marine boundary layer (MBL) temperature of 293 K, the rate coefficient is only 8.5 × 10 −16 cm 3 s −1 .…”

Section: Electronic Structure Calculations On Halogen Oxide Reactionssupporting

confidence: 86%

“…Other candidates besides OH acting as SO 2 oxidants might be halogen oxide radicals; however, to our knowledge respective rate constants are available in the literature only for the reactions of IO and ClO with SO 2 (Larin et al, 2000;DeMore et al, 1997), which are 3 and 6 orders of magnitude smaller compared to k SO 2 +OH , respectively. We have made ab initio transition state energy calculations for the reactions of SO 2 with ClO, BrO, IO, and OIO using quantum theory.…”

Section: Electronic Structure Calculations On Halogen Oxide Reactionsmentioning

confidence: 99%

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Missing SO<sub>2</sub> oxidant in the coastal atmosphere? – observations from high-resolution measurements of OH and atmospheric sulfur compounds

et al. 2014

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Abstract. Diurnal and seasonal variations of gaseous sulfuric acid (H 2 SO 4 ) and methane sulfonic acid (MSA) were measured in NE Atlantic air at the Mace Head atmospheric research station during the years 2010 and 2011. The measurements utilized selected-ion chemical ionization mass spectrometry (SI/CIMS) with a detection limit for both compounds of 4.3 × 10 4 cm −3 at 5 min signal integration. The H 2 SO 4 and MSA gas-phase concentrations were analyzed in conjunction with the condensational sink for both compounds derived from 3 nm to 10 µm (aerodynamic diameter) aerosol size distributions. Accommodation coefficients of 1.0 for H 2 SO 4 and 0.12 for MSA were assumed, leading to estimated atmospheric lifetimes on the order of 7 and 25 min, respectively. With the SI/CIMS instrument in OH measurement mode alternating between OH signal and background (non-OH) signal, evidence was obtained for the presence of one or more unknown oxidants of SO 2 in addition to OH. Depending on the nature of the oxidant(s), its ambient concentration may be enhanced in the CIMS inlet system by additional production. The apparent unknown SO 2 oxidant was additionally confirmed by direct measurements of SO 2 in conjunction with calculated H 2 SO 4 concentrations. The calculated H 2 SO 4 concentrations were consistently lower than the measured concentrations by a factor of 4.7 ± 2.4 when considering the oxidation of SO 2 by OH as the only source of H 2 SO 4 . Both the OH and the background signal were also observed to increase significantly during daytime aerosol nucleation events, independent of the ozone photolysis frequency, J (O 1 D), and were followed by peaks in both H 2 SO 4 and MSA concentrations. This suggests a strong relation between the unknown oxidant(s), OH chemistry, and the atmospheric photolysis and photooxidation of biogenic iodine compounds. As to the identity of the atmospheric SO 2 oxidant(s), we have been able to exclude ClO, BrO, IO, and OIO as possible candidates based on ab initio calculations. Nevertheless, IO could contribute significantly to the observed CIMS background signal. A detailed analysis of this CIMS background signal in context with recently published kinetic data currently suggests that Criegee intermediates (CIs) produced from ozonolysis of alkenes play no significant role for SO 2 oxidation in the marine atmosphere at Mace Head. On the other hand, SO 2 oxidation by small CIs such as CH 2 OO produced photolytically or possibly in the photochemical degradation of methane is consistent with our observations. In addition, H 2 SO 4 formation from dimethyl sulfide oxidation via SO 3 as an intermediate instead of SO 2 also appears to be a viable explanation. Both pathways need to be further explored.

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Section: Electronic Structure Calculations On Halogen Oxide Reactionssupporting

confidence: 86%

Section: Electronic Structure Calculations On Halogen Oxide Reactionsmentioning

confidence: 99%

Missing SO<sub>2</sub> oxidant in the coastal atmosphere? – observations from high-resolution measurements of OH and atmospheric sulfur compounds

et al. 2014

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“…A modelling study 29 found better agreement with field data when using the lower value for k 1 . Several later experimental studies confirmed the lower value of k 1 6,[30][31][32][33] and the current IUPAC room temperature value is (1.3 AE 0.2) Â 10 À14 cm 3 molecule À1 s À1 . 13 The only one of these studies to investigate the pressure dependence of k 1 was that of Daykin and Wine 28 who found no change in the rate coefficient over a pressure range of 40-300 Torr (M ¼ N 2 /O 2 /air).…”

Section: Introductionmentioning

confidence: 59%

Kinetics study of the reaction of iodine monoxide radicals with dimethyl sulfide

Gravestock

Blitz

Heard

2005

Phys. Chem. Chem. Phys.

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The temperature and pressure dependence of the rate coefficient for the reaction of iodine monoxide radicals with dimethyl sulfide (DMS), IO + DMS --> I + DMSO (1), was studied using laser induced fluorescence (LIF) to monitor the temporal profile of IO following 351 nm photolysis of RI/DMS/NO2/He (RI = CH3I/CF3I) mixtures. The study was performed over the range T = 296-468 K yielding a positive activation energy and k1 = (9.6 +/- 8.8) x 10(12) exp{-(1816 +/- 397)/T}. No dependence was observed on total pressure between 5-300 Torr. The rate coefficient at 296 K was determined as (2.0 +/- (0.6)(0.4)) x 10(-14) cm3 molecule(-1) s(-1), more than an order of magnitude smaller than a recent study but in reasonable agreement with the previous literature.

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“…CH 3 S(O)CH 3 was detected as a product in all three experiments, and Barnes et al calculated a semi-quantitative yield of (0.84 AE 0.4). Larin et al 19 measured an upper limit of k 1 (298 K) o 2.5 Â 10 À14 cm 3 molecule À1 s À1 by discharge-flow/ resonance-fluorescence detection of I-atoms. The current IU-PAC recommendation, 14 of k 1 (298 K) ¼ (1.5 AE 0.2) Â 10 À14 cm 3 molecule À1 s À1 is based upon the results of the three DF-MS studies, and the most recent publication, Gravestock et al 20 who used a pulsed laser photolysis technique to obtain k 1 (298 K) ¼ (2.0 þ0.4 À0.6 ) Â 10 À14 cm 3 molecule À1 s À1 , independent of bath gas pressure (5 o P He /Torr o 300).…”

Section: Introductionmentioning

confidence: 99%

The reaction of IO with CH3SCH3: products and temperature dependent rate coefficients by laser induced fluorescence

Dillon

Karunanandan

Crowley

2006

Phys. Chem. Chem. Phys.

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The technique of pulsed laser photolysis was coupled to laser induced fluorescence detection of iodine oxide (IO) to measure rate coefficients, k(1)(T), for the title reaction IO + CH3SCH3 --> products (R1). A value of k1(298 K) = (1.44 +/- 0.15) x 10(-14) cm3 molecule(-1) s(-1) was obtained, independent of bath gas pressure (50 < P((N2 or air))/Torr < 300). The expression k1(T) = (3.2 +/- 1.4)x 10(-13)exp[(-925 +/- 136)/T)] adequately described the data over the range of temperatures (256 < T/K < 341) covered. Uncertainties (2sigma) in the 298 K rate coefficient and the pre-exponential factor include an estimate of systematic error. The conventional Arrhenius behaviour of k1(T) and the lack of pressure dependence are suggestive of an abstraction mechanism, characterised by an energy barrier of E approximately 8 kJ mol(-1). The product yield for production of I-atoms was determined indirectly to be close to unity, indicating that the reaction proceeds via transfer of the O-atom from IO to CH3SCH3 to form CH3S(O)CH3. In general, the values of k1(T) measured in this work indicate that has little impact on the chemistry of the atmosphere.

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Measurement of the rate constants for the reactions of the IO• radical with sulfur-containing compounds H2S, (CH3)2S, and SO2

Cited by 15 publications

References 8 publications

Missing SO<sub>2</sub> oxidant in the coastal atmosphere? – observations from high-resolution measurements of OH and atmospheric sulfur compounds

Missing SO<sub>2</sub> oxidant in the coastal atmosphere? – observations from high-resolution measurements of OH and atmospheric sulfur compounds

Kinetics study of the reaction of iodine monoxide radicals with dimethyl sulfide

The reaction of IO with CH3SCH3: products and temperature dependent rate coefficients by laser induced fluorescence

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Measurement of the rate constants for the reactions of the IO• radical with sulfur-containing compounds H2S, (CH3)2S, and SO2

Cited by 15 publications

References 8 publications

Missing SO&lt;sub&gt;2&lt;/sub&gt; oxidant in the coastal atmosphere? – observations from high-resolution measurements of OH and atmospheric sulfur compounds

Missing SO&lt;sub&gt;2&lt;/sub&gt; oxidant in the coastal atmosphere? – observations from high-resolution measurements of OH and atmospheric sulfur compounds

Kinetics study of the reaction of iodine monoxide radicals with dimethyl sulfide

The reaction of IO with CH3SCH3: products and temperature dependent rate coefficients by laser induced fluorescence

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Missing SO<sub>2</sub> oxidant in the coastal atmosphere? – observations from high-resolution measurements of OH and atmospheric sulfur compounds

Missing SO<sub>2</sub> oxidant in the coastal atmosphere? – observations from high-resolution measurements of OH and atmospheric sulfur compounds