Environ. Sci. Technol. 1992,26,527-532
provide more insight into the origin of indoor aerosols. Acknowledgments We thank :Drs. L. Sheldon and T. Hartwell of Research Triangle Institute for their considerable cooperation in providing rapid access to the particulate filter samples and collected field data. We also acknowledge Drs. P. B. Ryan and J. D. Spengler for their instructive comments and Dr. Joseph Rizzuto for his contributions to the field study.
Registry No. Si, 7440-21-3; S, 7704-34-9; K, 7440-09-7; Ca, 7440-70-2; V, 7440-62-2; Mn, 7439-96-5; Fe, 7439-89-6; Ni, 744002-0; Cu, 7440-50-8; Zn, 7440-66-6; As, 7440-38-2;Se, 7782-49-2; Cd, 7440-43-9;Pb, 7439-92-1.
Literature Cited Marple, V. A.; Rubow, K. L.; Turner, W. A.; Spengler, J. D. J.-Air Pollut. Control Assoc. 1987, 37, 1303-1307. Leaderer, B. P.; Koutrakis, P. E.; Briggs, S. L. K.; Rizzuto, J. Atmos. Environ., in press. Cote, E. A.; Dietz, R. N. Environ. Znt. 1982, 8 , 419-433. Sinclair, J. D.; Psota-Keety, L. A.; Weschler, C. J. Atmos. Environ. 1988, 22, 461-469. Dockery, D. W.; Spengler, J. D. Atmos. Environ. 1981,15,
Glossary C
B Ci C O
E
Er Fd
F0"t
k P
&kerosene &is Qos &other &smoke Qwood Skerosene Ssmoke Swood ud
V
W
a U
number of cigarettes smoked slope of regression Ci on Co indoor aerosol concentration outdoor aerosol concentration home floor surface room floor surface particle deposition flux aerosol transport from indoors to outdoors hours of kerosene burning particle penetration efficiency flux of kerosene burning flux of particles from indoors flux of particles from outdoors flux of unknown indoor sources flux of cigarette smoking flux of wood burning particle emission rate of kerosene burning particle emission rate of cigarette particle emission rate of wood burning deposition velocity house volume hours of wood burning air exchange rate surface-to-volume ratio
335-341.
Lewis, C. W. J. Exposure Anal. Environ. Epidemiol. 1991, 1 , 31-44.
Leaderer, B. P.; Hammond, S. K. Environ. Sci. Technol. 1991,25, 770-777.
Pacyna, J. M. Norwegian Institute of Air Research, NILV Technical Report 10/82; 1983. Lebret, E.; McCarthy, J.; Spengler, J. D.; Chang, B. H. The Fourth International Conference on Indoor Air Quality and Climate, West Berlin, Germany, August 17-21, 1987. Colome, S. D.; Spengler, J. D.; McCarthy, S. Environ. Znt. 1982,8,197-212.
Koutrakis, P.; Spengler,J. D.; Chang, B. H.; Ozkaynak, H. Nucl. Znstrum. Methods Phys. Res. 1987, B22, 331-336. Santanam, S.; Spengler,J. D. 82nd Annual Meeting of Air and Waste Management Association, Anaheim, CA, June 25-30, 1989.
Received for review December 26, 1990. Revised manuscript received June 3, 1991. Accepted October 28, 1991. This work is supported by EPA Cooperative Agreement CR-814150. The field portion of this study was conducted by the Research Triangle Institute f o r the New York State Energy Research Development Authority (NYSERDA).
Oxidation of Alkyl Sulfides by Aqueous Peroxymonosulfate Eric A. Betterton Department of Atmospheric Sciences, University of Arizona, Tucson, Arizona 85721
The kinetics and mechanism of the oxidation of the alkyl sulfides dimethyl sulfide (DMS) and diethyl sulfide (DES) by peroxymonosulfate (HS05-)in aqueous solution were studied as a function of pH, temperature, and ionic strength. The rate law was found to be -d[R,S]/dt = kl[R2S][HS05-] + k2[R2S][S0,2-],where kl(DMS) = 0.13 f 0.09 M-' s-l, k2(DMS) = 29.8 f 0.08 M-' s-', kl(DES) = 0.29 f 0.01 M-l s-l, and k2(DES) = 11.1f 1.0 M-' s-l (25 "C; ionic strength, 0.2 M). The activation parameters for DMS and DES, respectively, are M * k , = 51.3 f 0.5 and 40.7 f 5.9 kJ mol-', AS*k, = -87.0 f 0.1 and -119 f 20 J K-' mol-', M * k 2 = 43.5 f 1.9 and 18.5 f 3.3 kJ mol-l, and AS*k2 = -70 f 7 and -164 f 18 J K-' mol-l. Dimethyl sulfoxide, dimethyl sulfone, diethyl sulfoxide, and diethyl sulfone were identified as reaction products. A comparison of the relative rates of oxidation of DMS and DES by HS05-, H202,and O3shows that HS05- may be a more important aqueous sink for the alkyl sulfides than H202. O3is, however, likely to be the dominant aqueous oxidant under conditions that could be expected in remote marine clouds. Introduction There is a need to develop a sound understanding of the kinetics and mechanism of the atmospheric chemical 0013-936X/92/0926-0527$03.00/0
transformations of reduced sulfur compounds including alkyl sulfides such as dimethyl sulfide (DMS) and diethyl sulfide (DES). Global source strengths of DMS are estimated to be on the order of (40-45) X 10l2g of S year-l (I),with the major source being the enzymatic decomposition of dimethylsulfonium propionate, an intracellular osmoregulatory substance in marine phytoplankton (2). Sources of DES include animal waste, natural gas, and wood pulping (3). The ultimate product of most sulfur oxidation processes is particulate sulfate, a key component of cloud condensation nuclei, suggesting a possible link between DMS emissions and cloud microphysics (4, 5). The relationship between DMS and sulfate levels seems to be nonlinear however (6). DMS is also oxidized to methanesulfonic acid, which recent studies suggest is metabolized to carbon dioxide and sulfate by certain soil bacteria (7). A detailed understanding of the cycling of sulfur in the environment is therefore crucial to our understanding of cloud microphysics, global albedo, and climate change phenomena. It is thought that the dominant loss process for most of the atmospheric reduced sulfur compounds involves gasphase oxidation by OH', during the day, and oxidation by at night (8,9).For example, the lifetime, 7 (defined as 7 = l/k[oxidant]), for loss of DMS by OH', is approx-
0 1992 American Chemical Society
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imately 48 h when [OH',] = lo6 molecule ~ m -while ~ , the expected lifetime for loss by NO3', is 2 h when [NO;,] = 5 pptv (lo8 molecule ~ m - ~ If) .the nighttime NO3*,level is closer to 0.04 pptv (lo6 molecule ~ m - ~as) ,Toon et al. (10) have suggested, then the lifetime of DMS with respect to this sink would rise to approximately 400 h. Field studies (11) indicate that the residence time of DMS in the marine boundary layer may be as low as 0.5-3 h. Oxidation by O3 g is too slow to be of importance, and organic sulfides are not subject to significant photolysis under tropospheric conditions (8). Most attention has been given to studies of the gas-phase reactions of biogenic sulfides such as H2S, (CH3),S, CH3SH, CH3SSCH3,CS2, and COS with OH',, and, to a lesser extent, with N03*gand IO', (1.0, but comparatively few aqueous-phase loss processes have so far been investigated. This is partly because the reduced sulfur compounds have moderately low Henry's law coefficients, and they are therefore likely to exist in the aqueous phase only at relatively low concentrations, and also because few strong aqueous oxidants that react rapidly with sulfides have so far been identified at high concentrations. Adewuyi (12) has recently reviewed the oxidation of biogenic sulfur compounds in aqueous solution. Potential oxidants for DMS include OH', 0,, H202, and 0, (in the presence of catalysts and/or light). The rate of reaction with OH' is almost diffusion controlled (5.2 X lo9M-l s-l) (13),and aqueous ozone also reacts rapidly with alkyl sulfides (lz > 2 X 10-j M-l s-l) (14). The rate of reaction of DMS with H,02 is acid catalyzed a t pH 6. In the acid-independent region, the second-order rate constant is 3.4 X lo-, M-l s-l (12). The autoxidation reaction (with 0,) is slow in the absence of trace-metal catalysts (2.2 X s-l). Photosensitizers, including humic acid, accelerate the rate of photooxidation: Brimblecombe and Shooter (15) reported that the firstorder rate constant for photooxidation of DMS exposed to sunlight in seawater containing natural photosensitizers is 2.4 X s-l. While peroxymonosulfate (HS05-) has not yet been detected in hydrometeors, Jacob's (16) cloudwater chemistry model predicts that HS05- may occur in remote marine and continental clouds a t levels approaching 2.5 pM. It is doubtful if HSOc can compete effectively with OH', during the day because the low Henry's law constant of the reduced sulfur compounds limits the total S(-11) load in cloudwater, but HSO, could nevertheless be a potentially important aqueous sink for DMS and other reduced sulfur compounds. Peroxymonosulfate, a monosubstituted derivative of H202,is a powerful oxidizing agent (17) (EHso6p- .= 1.82 V). Peroxymonosulfurous acid (H2S05), w ich is also known as Caro's acid, is a strong acid with pKal < 0 and pK, = 9.88 f 0.1 (15 "C) (18). Therefore HSO; (and not H2S0, or SO:-) is usually the only significant species over the pH range of most atmospheric studies. HSO, is stable in the acidic to neutral pH range and also at high pH, but it decomposes fairly rapidly to yield oxygen when pH = pKa2(19). This broad range of stability makes it possible for HS05- to persist in remote clouds after all available reductants have been depleted. The Henry's law constants for DMS and DES are 0.56 and 0.46 M atm-l (25 "C), respectively (20, 21). The magnitude of these constants implies that for typical cloud droplets of 10-pm radius the characteristic times for gasphase diffusion to the droplet, for achieving interfacial equilibrium, and for aqueous-phase diffusion in the droplet are all on the order of a few milliseconds or less (22). 528
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Therefore it is the chemical reaction time that will determine the overall rate of oxidation of S(-II). Further, since chemical reaction times are expected to be on the order of hours to days under natural conditions, we can assume an open system with a constant supply of sulfur gases. Our previous work has demonstrated that HSO, rapidly oxidizes SOzaq(23) and H2S, (24). Here we report a study of the kinetics and mechanism of the oxidation of DMS and DES by HS0,- in aqueous solution as a function of pH, temperature, and ionic strength. The results may be useful for modeling natural biogeochemical cycles and may also be of interest in assessing the suitability of peroxymonosulfate for treating industrial waste streams containing malodorous organic sulfides. Experimental Section
AR grade reagents were used whenever possible. Important exceptions were dimethyl sulfide (Aldrich, anhydrous, 99+%), diethyl sulfide (Kodak), and peroxymonosulfate (Aldrich Oxone, 2KHS05.KHS04-K2S04). Distilled water was further purified by passing it through a Millipore MilliQ water system (18 MQ cm resistivity). Buffers with an ionic strength of 0.2 M were prepared from NaH2P04/Na2HP04(pH 6-49, NaHC03/Na2C03 (pH 9-11), and NaC104/NaOH (pH 12-13). They were filtered through 0.4-pm Nucleopore filters and deoxygenated by sparging with nitrogen immediately before use. DMS and DES solutions were prepared by dissolving the calculated volume of reagent in the appropriate buffer [p:' = 0.846 g mL-l (DMS);p i o = 0.837 g mI-l (DES)] (25). The final concentration, after mixing with an equal volume of HS06- in a cuvet, was 1 X M. DMS is subject to only very slow autoxidation in the absence of catalysts or photosensitizers (12). Peroxymonosulfate solutions were prepared by dissolving the calculated mass of the triple salt Oxone (2KHSO5.KHSO4.K2SO4)in the appropriate buffer and readjusting the pH if necessary before diluting to the mark. The final concentration after mixing in the cuvet was 1 X or 2 x M HSOS-T, i.e., at least a 10-fold excess over the alkyl sulfide concentration (HSOcT = HSO,- + Sob2-). Peroxymonosulfate solutions were prepared quickly and used immediately because of the possibility of decomposition particularly when pH pKa2. Spectrophotometric measurements were made with a Hitachi U 2000 spectrophotometer fitted with a thermostated sample compartment. Reservoirs of the reactants to be pumped through a flow-through cuvet were held at the required temperature in a water bath. The pH was measured with an Orion SA 720 pH meter and combination glass electrode calibrated with Radiometer buffers. Glassware was soaked in a cleaning solution of ammonium peroxydisulfate in concentrated H2S04to minimize the risk of trace-metal contamination. Reagents were protected from light. Although neither DMS nor DES is highly soluble in water (17.7 and 3.1 g/L, respectively) (26),they each have a fairly strong absorption band near 200 nm: it was found that for DMS, czo0 = 9.4 x lo2 M-l cm-l; and for DES, €207 = 1.7 x lo3 M-l cm-I. Beer's law was found to be obeyed up to a t least 2 absorbance units. The kinetics were studied by monitoring the rate of loss of this band. In order to improve the signal to noise ratio, the wavelength was set to 220-240 nm, depending on the pH. This wavelength region also corresponds to the region of maximum absorbance difference between the reactant sulfide and product sulfoxide. HS05- did not interfere since it shows only weak absorption in this region (ez7,, = 5.5 M-' cm-l) (18) and it was present at pseudoconstant concen-
-
L *O
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v
C -
-2
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101 5
I
-3
200
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400
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800
Time (s)
7
9
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13
PH
Figure 2. Plot of second-order rate constant, k,,,l[HSO,-],, vs pH for the reaction of peroxymonosulfate with dimethyl sulfide and diethyl sulfide; 25 O C , ionic strength, 0.2 M. 1
-5
'
0
I
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1
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Figure 1. Typical plots obtained for the oxidation of the dimethyl sulfide by peroxymonosulfate (a) at low pH (2 X lo-' M [HSO,-],,) and (b) at high pH (1 X lo-' M [HSO,-ITo). The error bars show the range for three replicate experiments and the solid lines are the linear least-squares fits to the median data points.
tration. However, SO2- absorbs more strongly (ezTo = 85.5 M-l cm-') (18)and the background absorbance therefore increased with pH. The half-life varied from approximately 3 to 300 s depending on the conditions and this necessitated three approaches to mixing the reagents. Slow reactions were initiated by manually injecting 1mL of each reactant into a 1-cm quartz cuvet (Hellma), while faster reactions were initiated by mixing equal volumes of the reactants through a polypropylene T-piece into a 1-cm quartz flow-through cuvet (Hellma) using a dual-head peristaltic pump (Cole Parmer) fitted with silicone tubing. To begin a measurement the pump was stopped abruptly and absorption measurements commenced immediately. The dead time between mixing and observation was calculated to be > Ka2,i.e., when pH 5 ~ 7 . 9 kobsd/ [HS05-lT k1 + Ka2kZ/ [H+1 (6) so that a linear plot of k&sd/ [HSO5-ITvs 1/ [H+] gives k, (intercept) and k 2 (slope/Ka,). Similarly, when [H+] 2 X 23 day 0.0778 9h
lifetimeb H202
O3
116 day =50 s 136 day
