Rate of Reaction of Methanesulfonic Acid, Dimethyl Sulfoxide, and

Chapter 33

Rate of Reaction of Methanesulfonic Acid, Dimethyl Sulfoxide, and Dimethyl Sulfone with Hydroxyl Radical in Aqueous Solution

Downloaded by UNIV QUEENSLAND on May 5, 2013 | http://pubs.acs.org Publication Date: April 27, 1989 | doi: 10.1021/bk-1989-0393.ch033

P. J. Milne, Rod G. Zika, and Eric S. Saltzman Division of Marine and Atmospheric Chemistry, Rosenstiel School of Marine and Atmospheric Science, University of Miami, Miami, FL 33149-1098 Methanesulfonic acid, dimethyl sulfoxide and dimethyl sulfone are potential intermediates in the gas phase oxidation of dimethylsulfide in the atmosphere. We have measured the rate of reaction of MSA with O H in aqueous solution using laser flash photolysis of dilute hydrogen peroxide solutions as a source of hydroxyl radicals, and using competition kinetics with thiocyanate as the reference solute. The rate of the reaction k'(OH + SCN ) was remeasured to be 9.60 ± 1.12 x 10 M s , in reasonable agreement with recent literature determinations. The rates of reaction of the hydroxyl radical with the organosulfur compounds were found to decrease in the order DMSO (k' = 5.4 ± 0.3 x 10 M s- ) > MSA (k' = 4.7 ± 0.9 x 10 M s ) > D M S O (k' = 2.7 ± .15 x 10 M s ). The implications of the rate constant for the fate of MSA i n atmospheric water are discussed. -

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Organosulfur compounds, of which dimethylsulfide (DMS) appears to be the most abundant, are the principal precursors of sulfur dioxide and non sea-salt sulfate aerosol in non-polluted air, and have a major impact on the global tropospheric sulfur cycle. The primary mechanism for atmospheric oxidation of these compounds is gas phase reaction with O H or possibly ΝΟ3· radicals. Until quite recently (1) the oxidation product of organosulfur gases in the atmosphere was thought to be primarily sulfur dioxide, which would then be subsequently oxidized, again quantitatively, to sulfate. However, it is now apparent that a number of other intermediate products and chemically reactive transients may also be formed from organosulfur compounds in passing from the (-II) to the (+VI) oxidation state. This work examines the possible further oxidation of one of the more stable oxidation products, methanesulfonic acid, by hydroxyl radicals. Methanesulfonic acid, although it comprises a relatively small fraction of total non sea-salt aerosol sulfur, has been shown (2) to be a ubiquitous component of marine aerosols. Its occurrence and distribution have been suggested as of use as an in situ tracer (3.4) for oceanic emissions and subsequent reaction and deposition pathways of organosulfur compounds and dimethyl sulfide in particular. 0097-6156/89/0393-0518S06.00/0 * 1989 American Chemical Society

In Biogenic Sulfur in the Environment; Saltzman, E., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 1989.

33.

MILNE ETAL.

Rate ofReaction ofHydroxyl Radical in Aqueous Solution 519

Specifically we wished to measure the rate of reaction of O H with MSA to enable modelling calculations of the stability of MSA in aerosol droplets. The one reported measurement of this rate (5), using pulse radiolysis techniques, 3.2 x 10 M * s" , is fast enough to suggest that this reaction pathway could be an important sink for MSA. This is of interest in explaining an apparent discrepancy that exists between laboratory and field studies of the oxidation of dimetnyl sulfide. Although a number of laboratory studies (6-9) show that MSA is the major stable product, and S 0 a minor one, field observation suggest MSA is only a minor (10%) fraction (2) of total non-sea-salt sulfur in marine aerosols. Two possible rationalizations of this are that i) MSA is subject to further reaction in marine aerosols and ii) other reaction pathways of dimethyl sulfide, or perhaps other non-methylated sulfur compounds should be considered. Dimethyl sulfoxide, (DMSO) and dimethyl sulfone, (DMS0 ) have been reported in rain water samples (1Q). As is the case for MSA, the vapour pressures of these compounds are such that they are much more likely to be partitioned into heterogenous (aqueous) phases than to remain in the gas phase. A second aim of this work was to observe possible reaction pathways of these organosulfur compounds, which are also potentially stable oxidation products of dimethyl sulfide. 9

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Methods and Experimental Approach The experimental approach used in this study combined time resolved laser flash photolysis with competition kinetics. Hydroxyl radicals were produced in solution by the laser flash photolysis of dilute solutions (5 - 50mM) of hydrogen peroxide which undergoes the homolytic cleavage H 0 2

2

> 20H hi/

The excitation source for this step was an excimer laser (EMG 201MSC, Lambda Physik, Acton Ma) operated at the Kr-F (248nm) line. The laser produced a ~25ns width pulse of rectangular cross section, a portion of which was allowed to impinge on a (1cm) quartz flow cell. Test solutions were pumped through the cell by a peristaltic pump connected to an external reservoir, whereby solutions could be exchanged or reagent composition readily adjusted. Absorbance was monitored at right angles to the pump beam with a xenon flash lamp (FX 193 U , E G & G Electro-Optics, Salem, MA) collimated to pass through the irradiated portion of the sample cell and collected with a monochromator (HR-320,Instruments SA, Metuchen, N J ) acting as a spectrograph for an intensified gated diode array (DARSS, TN-6133, Tracor Northern). The diode array was interfaced through an optical multichannel analyzer (Tracor Northern TN-1710) to a laboratory computer (Hewlett Packard HP-85). The timing sequence of the experiment with respect to i) the pump laser flash, ii) the probe beam lamp flash, and iii) the gate pulse on the diode array detector was set with the use of a digitial delay/pulse generator (DG535, Stanford Research Systems, Palo Alto, Ca). In this way, spectra of approximately 300nm width were obtained at 40ns intervals at selected times after to, as defined by the laser flash. In the experiments performed here, the region 275-575nm was monitored at intervals up to 2^sec after the photolyzing flash. The experimental system used in the time resolved absorption measurements is outlined in Figure 1. Analytical grade potassium thiocyanate (Fisher), hydrogen peroxide (Mallinckrodt), dimethyl sulfoxide, (Fluka), dimethyl sulfone (Sigma) were used In Biogenic Sulfur in the Environment; Saltzman, E., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 1989.

520

BIOGENIC SULFUR IN THE ENVIRONMENT


L

L

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- H t a

OMA

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C

MC

DO

i

EL

Figure 1. Experimental system used in the time resolved absorption measurements. (EL=excimer laser, KrF, 248nm: D G = delay generator: OMA=optical multichannel analyser: MC=monochromator and gated diode array detector: C=cell: X=xenonflashlamp: L=lenses )


33. MILNE ETAL.

Rate of Reaction ofHydroxyl Radical in Aqueous Solution 521

as received. Methanesulfonic acid (Morton Thiokol) was specified as high purity (99.5%), yet two batches had a light brown colouration. L C and NMR examination of neat samples of the methanesulfonic acid did not however reveal any impurity.

Competition Kinetics Gas phase kinetic studies of the reactions of hydroxyl radical are most conveniently carried out with direct monitoring of the O H radical with time using laser induced fluorescence (H). The low absorption coefficient of the aqueous hydroxyl radical (€igg - 540 M " cnr , (12)) precluded the direct measurement of this reactant species by its absorbance. Also, the absence of a readily observable product species for the reaction of O H + MSA at the wavelength range (2/5-575 nm) easily accessible in our experiments, has lead us to monitor the concentration of O H in solution indirectly by competition kinetics (13), measuring the absorption of the thiocyanate radical anion (c^gQnm = 7600NPcm-iQ2)). Competition kinetics uses the absorption of a transient (or stable) product formed from the reaction of O H with a reference solute. This absorption will be suppressed in the presence of a competitor solute, as O H radicals will react with either the reference or competitor solutes in proportion to the products of the concentration and rate constants of the two competing reactions. Specifically, for the reactions: 1

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nm

OH

+ SCN-

OH +

X

— > absorbing

product

— > non-absorbing

product

it follows that k(OH+X) k(OH+SCN-)

=

A (Ao-A)

[X] [SCN-]

(1)

Here k(OH+X) is the absolute rate of O H with competitor X , k(OH+SCN) the absolute rate of O H with reference solute, SCN-, AQ is the absorbance of the reference product when X is absent, and A is the absorbance in the presence of some concentration of X . Significant assumptions of this scheme are: i) the kinetics can be described as simple bimolecular reactions ii) there are no other absorbing species at the wavelength being measured iii) there are no secondary reactions, of the reference or the competitor, occurring in the time frame being studied. Rate Qf Formation of (SCN) 2

Hydroxyl radical oxidation of thiocyanate ion in acid and neutral solution leads to the formation of a transient species which absorbs strongly at a wavelength of 475nm. Figure 2 shows the growth in absorption of this species at intervals from 20ns to 800ns. Historically, this absorption was first attributed (14) to the SCN radical formed by O H + SCN- —> SCN + OH-


In Biogenic Sulfur in the Environment; Saltzman, E., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 1989. 2

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Figure 2. Growth of the absorption due to the formation of ( S C N V at intervals ranging from 40ns to 800ns after the laser pulse. (KrF laser, excitation wavelength 248nm, [H 0 ] = 50mM, [SCN*] = 0.5mM)


2 » S

M

33. MILNE ETAL.

Rate ofReaction ofHydroxyl Radical in Aqueous Solution 523

The data was also shown (i£) to be consistent with the transient species being (SCN) ", analogous to the dihalide anion radicals, produced by 2

SCN + SCN-—> (SCN)

2

Further investigations (16) have suggested an even more complex description of the mechanism: O H + SCN- — > SCNOHSCNOH- —> SCN + OHDownloaded by UNIV QUEENSLAND on May 5, 2013 | http://pubs.acs.org Publication Date: April 27, 1989 | doi: 10.1021/bk-1989-0393.ch033

SCN + SCN- —> (SCN)

2

SCNOH- + SCN- — > (SCN)

2

+ OH-

Correspondingly, some uncertainty has existed as to the value of the rate constant for the reaction of O H + SCN", depending upon the kinetic scheme used to fit the experimental data. Due to the large number of pulse radiolysis competiton kinetic studies performed using thiocyanate as the reference system, the absolute value of this rate is of some concern and has been in fact arbitrarily adjusted (1.1 x 10 M * s , (12)) to accord with other reference solutes. Independently, a very similar value (k* = 1.08 ± 0.10 x 10 M s ) was measured (12). 10

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Results and Discussion Figure 3 shows a plot of the rate of formation of the (SCN) - in our experiments measured at two different thiocyanate concentrations. Linear fits of these and other data enabled evaluation of the slopes of lnKA^-AVAnJ, where A is the maximum absorbance. Attempts were made to fit the kinetic data with the reaction scheme suggested by Ellison et al. (12), but the lack of sufficient data points in the time domain, together with the inherently ill-conditioned nature of a triple exponentialfit,meant that this was not meaningful for our data set. Figure 4 shows the observed k* values plotted as a function of the thiocyanate concentration, the slope of which yields a value of 9.60 ± 1.12 x 10 M s" . This value is in fair agreement with an accepted value of 1.1 x 10 M " 2

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s- (12,12). Competition experiments in which increasing amounts of the substrate of interest were addea to a known concentration of thiocyanate and peroxide before photolysis were carried out for MSA, DMSO, and DMSCK Typical conditions for these experiments were [SCN-] = 0.5mM, [ H 0 ] = 50 mM, [MSA] = 0~75mM, [DMSO] = 0~lmM, and [DMSOJ = 0~150mM. Solutions were not degassed, although an experiment with DMSO which had been sparged and blanketed with nitrogen gave a rate constant similar to those that had not. Solutions were also unbuffered, although their pH was monitored (~ pH 2 for experiments with MSA, pH 4-5 otherwise). In none of the experiments was there indication of absorbance by species other than the thiocyanate radical, which supports the assumption that any reactions between uncomplexed SCN radicals and other species present are unimportant. Figures 5a),b) show the diminution of the thiocyanate radical absorption at 480 nm (A) as increasing amounts of each added competitor substance were added for the substrates added. Using equation (1), plots of A Q / A VS [S]/[SCN-] were made, the slopes of which yielded k(OH + s)/k(OH + S C N ) , where s is 2


2

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BIOGENIC SULFUR IN THE ENVIRONMENT

0-1-2-


Rate of Reaction of Methanesulfonic Acid, Dimethyl Sulfoxide, and

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