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J . Phys. Chem. 1992, 96, 3317-3320
Pulse Radiolysis Study of the Reaction of OH' Radicals with Methanesulfonate and Hydroxymethanesulfonate Terese M. Olson* Department of Civil Engineering, University of California, Irvine, California 9271 7
and Richard W. Fessenden Radiation Laboratory and Department of Chemistry and Biochemistry, University of Notre Dame, Notre Dame, Indiana 46556 (Received: September 24, 1991; In Final Form: December 6 , 1991)
Pulse radiolysis studies of the reaction kinetics of methanesulfonate and hydroxymethanesulfonate with hydroxyl radicals were conducted by relative rate methods. The rate constant for the reaction of CH3S03-with OH' was determined as 1.24 (f0.07) X lo7 M-' s-I at an ionic strength of 0.2 M, with thiocyanate as the reference solute. For the same ionic strength conditions, the rate constant for the attack of OH' with CH2(OH)S03-, k = 2.65 (k0.12) X lo* M-' s-l , w as more than an order of magnitude greater than the rate constant for methanesulfonate. Previous estimates of the rate constants for the methanesulfonate reaction varied by approximately 2 orders of magnitude; our results are in best agreement with the slowest value. Our rate constant for hydroxymethanesulfonate was also significantly smaller than two previously reported values. Both reactions are independent of pH over the pH range studied (3.8-6.5) and are nearly independent of ionic strength. These results imply that the oxidation of methanesulfonate by OH' is too slow to be significant in typical remote cloud droplets, but may be important in aerosols. The reaction of OH' with hydroxymethanesulfonate would be significant in atmospheric M. droplets when hydroxyl radical concentrations are at least
Introduction Alkanesulfonates form and are present in atmospheric droplets and aerosols; however, their fates are not well known. One factor contributing to this uncertainty is the lack of agreement between reported rates of their oxidation by hydroxyl radicals, a potentially important sink reaction. Methanesulfonic acid aerosol, for example, is readily produced upon the photolysis of reduced sulfur gases, such as dimethyl sulfide.1-2 Two reported rate constants for the reaction of CH3S03-with OH', however, vary by almost 2 orders of m a g n i t ~ d e . ~Consequently, .~ the predicted rates of this reaction for atmospheric conditions range from insignificant to significant, depending on which rate constant is assumed. Hydroxymethanesulfonate, CH2(0H)S03-, is another important alkanesulfonate that is formed in atmospheric droplets by the addition reaction of bisulfite and formaldehyde.s The formation of this adduct is particularly favored during nighttime periods, when photochemical oxidants are depleted, since sulfite oxidation by H202or O3is typically faster than the adduct formation rate.6 Once formed, hydroxymethanesulfonate does not appreciably react directly with ozone or peroxide. Its oxidation by hydroxyl radicals, however, has been proposed as a potential additional formation pathway for sulfate and acidity in The reaction kinetics of OH' with hydroxymethanesulfonate have also been studied by Martin et aL9 using Fenton's reagent. They determined a rate constant of 1.25 X lo9 M-' s-I, and concluded that the reaction represents a potentially important indirect oxidation pathway for SO2 in droplets. Their kinetic experiments were conducted in the presence of high concentrations of chloride ions, however, which are known to react relatively rapidly with OH'. Another lower limit for the rate constant was determined as 1 X lo9 M-l s-I by Deister et a1.I0 In this study, hydroxyl radicals were generated by the steady-state photolysis of nitrate ions. The lower limit for the rate constant was estimated on the basis of competition with benzene as the reference solute. A pulse radiolysis study of the reaction kinetics of methanesulfonate and hydroxymethanesulfonate with OH' was undertaken to resolve the question of their reactivity. While primary ionic strength effects are not expected to be significant for reactions involving the uncharged hydroxyl radical, the reactions could typically occur in strong electrolyte systems, such as aquated haze aerosols. The effects of ionic strength on the rates of these re-
* To whom correspondence should
be addressed.
0022-365419212096-3317$03.00/0
actions, therefore, were also investigated. Experimental Procedures Materials. Solutions were prepared from analytical grade reagents and water from a Millipore Milli-Q purification system. Prior to each experiment, the solutions were saturated with N20 to eliminate reactions with hydrated electrons and to purge oxygen from the sample. Commercial sources of sodium hydroxymethanesulfonate (Aldrich), sodium methanesulfonate (Aldrich), and KSCN (Fisher) were used without further purification. The pH and ionic strength were adjusted with H2S04(Fisher) and Na2S04 (Fisher), respectively. Methods. Pulse radiolysis was used to study the rates of the initial attack of hydroxyl radicals with the alkanesulfonates. The radiolysis apparatus consisted of a linear accelerator that produced 5-nspulses of 8-MeV electrons. Details of the facility are described elsewhere." An N20-saturated KSCN dosimeter solution was used to determine the average radiation dose prior to conducting the relative rate experiments. Only pulses within a narrow dose range from 8 to 11 Gy were utilized. Assuming the radiation chemical yield of OH', G(OH'), is approximately 6.0,12the initial concentration of OH' produced after the pulse was estimated as 6 pM. Relative rate techniques were employed to determine rate constants, since the reaction products could not be observed spectrophotometrically, and thiocyanate was chosen as the reference solute. The experiments were conducted in a flow-through, ~~~
~~
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( I ) Hatakeyama, S.; Akimoto, H. Geophys. Res. Len. 1982, 9,583-586. (2) Hatakeyama, S.; Akimoto, H. J. Phys. Chem. 1983,87, 2387-2395. (3) Milne, P. J.; Zika, R. G.; Saltzman, E. S. In Biogenic Suljur in rhe Enuironmenf;Saltzman, E. S . , Cooper, W. J., Eds.; ACS Symposium Series 393; American Chemical Society: Washington, DC, 1989; pp 518-528. (4) Lind, J.; Eriksen, T. E. Radiochem. Radioanal. L e f f . 1975, 21, 177-1 81. (5) Munger, J. W.; Tiller, C.; Hoffmann, M. R. Science 1986. 23/. 247-249. (6) Boyce, S. D.; Hoffmann, M. R. J. Phys. Chem. 1984,88,474&4146. (7) Jacob, D. J. J. Geophys. Res. 1986, 91, 9807-9826. (8) Olson, T. M.; Hoffmann, M. R. Amos. Enuiron. 1989.23.985-987, (9) Martin, L. R.; Easton, M. P.; Foster, J. W.; Hill, M. W. Afmos. Enuiron. 1989, 23, 563-568.
(IO) Deister, U.; Warneck, P.; Wurzinger, C. Eer. Eunrenges. Phys. Chem. 1990, 94, 594-599.
( I I ) Schuler, R. H. Chem. Ed. 1985, 2, 34-47. (12) Schuler, R. H.; Hartzell, A. L.; Behar, B. J. Phys. Chem. 1981, 85, 192-1 99.
0 1992 American Chemical Society
3318 The Journal of Physical Chemistry, Vol. 96, No. 8,1992
1 cm path length quartz cell, while N 2 0 was continually bubbled through the feed solution. The extent of the oxidation of SCNwas determined by spectrophotometric detection of the radical intermediate, (SCN)
