17, No. 3. pp. 625432,
000..5981/83/03062548 0 1983 Pergamon
STUDY OF NONAQUEOUS-PHASE OF SULFUR DIOXIDE
103.00/O Press Ltd.
B. D. HoLT, P. T. CUNNINGHAM, A. G. ENGELKEMEIR, D. G. GRACZYK and R. KUMAR Chemical Engineering Division, Argonne National Laboratory, Argonne, IL 60439, U.S.A. (Receioedfor publication 13 August 1982) Abstract-In a study of the mechanisms of atmospheric sulfate formation, oxygen isotope ratios were measured in sulfates and in the SO, and water vapors from which they were formed, in the absence of liquid water. In a 3-f glass chamber, SO2 and water vapor of various i*O contents were isotopically equilibrated, and then air oxidation of the SO1 to sulfate was performed by four different methods: high-voltage discharges, NOa addition, gamma irradiation and adsorption on activated charcoal. Isotopic equilibration between SO2 and water vapor proceeded rapidly, resulting in a strong dependence of the 6180 of the sulfate on that of the water vapor. Oxidation of SOa on dry charcoal occurred through the apparent formation of 9-oxygen, 2-sulfur, chemisorbed molecules which decomposed to sulfate in leach water. The hisO so:- vs 6’80*pg relationships observed for these four nonaqueous-phase oxidations of SOa to sulfate, together with those m three previously reported aqueous-phase oxidations ( Fe3+ -catalyzed air oxidation, charcoal-catalyzed air oxidation and Ha O2 oxidation), were compared to sulfate in rain and snow collected at Argonne, IL. The hi80 of sulfate in precipitation water was significantly higher than could be accounted for by any of the several oxidation reactions that were investigated as possible pathways in the formation of secondary sulfates in the atmosphere, either singly or in combination. 1. INTRODUCTION We are conducting an investigation into the various homogeneous and hetrogeneous mechanisms whereby gaseous SO, is oxidized to SO:-. Stable-oxygenisotope ratio ( “O/i6 0) measurements can be used to deduce the reaction sequence that results in the formation of a given sample of sulfate. This deduction is possible because different reaction sequences affect the oxygen isotope of the product sulfate differently, given the same starting reactants. This unique analytical tool has been developed in our laboratory and applied to atmospheric sulfates to derive their formation mechanisms. As previously reported (Holt et al., 1981a), the average seasonal 180 contents of the sulfates found in atmospheric precipitation at Argonne, IL, are significantly higher than can be fully accounted for by aqueous-phase mechanisms of transformation of SO, to sulfate by Fe3* -catalyzed air oxidation, charcoalcatalyzed air oxidation and H202 oxidation. (The isotopic results reported for these three aqueous-phase reactions are reproduced in Fig. 1 as curves 1,2 and 3, respectively.) In this paper we discuss the results from a set of reaction sequences that involve the oxidation of SOZ in the absence of liquid water. The isotopic equilibration between SO, and water vapor, somewhat comparable to that of the C02water vapor system (Bottinga and Craig, 1969), necessarily affects the oxygen isotopy of all nonliquid transformations of atmospheric SO, to sulfate. The dependence of S l8 0 [deviation in parts per thousand * Paper presented in part at the Fourth International Symposium on Stable Isotopes, Julich, F.R.G., 23-26 March, 1981.
( %,) the 18O/16O ratio of the sample from that of standard mean ocean water @MOW)] of SOZ on the 6180 of water vapor with which it is associated was, therefore, investigated as part of the study of sulfate formation in the absence of liquid water. In one set of oxidation experiments, a high-voltage spark was used to induce oxidation of SO, in humidified air, which simulated electrical activity in thunderstorms. Little is known about the involvement of lightning in the oxidation of atmospheric S02, possibly through the formation of oxides of nitrogen. In another set of experiments, sulfate was formed by the addition of NO, to SO, in humidified ait. In the atmosphere NO, results from combustion emissions, atmospheric electrical activity, and photochemical activity at high altitudes. Other workers have studied the oxidation of SO, by NO, with various temperatures, catalytic solid substrates, and concentrations of water vapor, SO2 and NO, (Stopperka et al., 1968; Kunin and Epifanov, 1968; Clarke and Britton, 1977; Barbaray et al., 1978; Schroeder and Urone, 1978; Schryer et al., 1979). Oxidation of SO, in humidified air was also induced by irradiation of the gas mixture by OH-radicalproducing gamma rays (Gordon and Mulac, 1975) somewhat in simulation of photochemical production of OH radicals in the atmosphere (Davis and Klauber, 1975; Wood et al., 1975; Castleman et al., 1975; Castleman and Tang, 1977). The oxygen isotopy of the transformation of SO, to sulfate on dry charcoal was investigated because of the probable similarity, if not identity, to the oxidation associated with adsorption on soot particles in the atmosphere (Novakov et al., 1974; Barbaray et al., 1977; Rosen and Novakov, 1978; Novakov, 1978; 625
B. D. HOLT et al.
10 u ;;; z
0 m %
Li Water (“IoO)
Fig. 1. Comparison of sulfates formed by seven laboratory methods of preparation and sulfates in precipitation water. Curve 1, aqueous air oxidation with Fe”+ catalyst; Curve 2, aqueous air oxidation with charcoal catalyst; Curve 3, aqueous H,O, oxidation; Curve 4, electric spark in humidified air; Curve 5, NO2 in humidified air; Curve 6, gamma irradiation in humidified air; Curve 7, water vapor, air, and SO, adsorbed
Toossi, 1978). Experimental conditions were varied to elucidate the mechanism of SO, oxidation on the “dry” charcoal (Pierce et al., 1951; Smith, 1959; Davtyan and Ovchinnivkova, 1955; Tartarelli et al., 1978). Reactions not isotopically examined were ultraviolet-induced oxidation of SO, in humidified air (Cox, 1973; Shen and Springer, 1976; Kasahara and Takahashi, 1976) and surface oxidations on non-sootlike particles (Liberti et al., 1978; Corn and Cheng, 1972; Meagher et al., 1978; Judeikis ef al., 1978; Haury et al., 1978). Mechanisms of formation of primary sulfates, i.e. sulfates formed in combustion sources before discharge to the atmosphere (Dietz, 1979; Dietz and Garber, 1978; Evans and Targett, 1978), are presently under investigation at our laboratory.
3 - LITER FLASK
of SO, and water mzpor
of SO2 were enclosed
in a sealed glass ampoule (10 cm x 7 mm OD) and placed in a 3.1-f reaction chamber (Fig. 2). The chamber without the platinum electrodes shown in Fig. 2) was then evacuated. Fifty microliters of degassed liquid water were cryogenically transferred into the reaction chamber. The water was completely evaporated in the chamber and room air was admitted through a train of purification reagents [Ascarite to remove CO,; Drierite and Mg(ClO,), to remove water vapor] to achieve atmospheric pressure. The reaction chamber was removed from the vacuum line and vigorously shaken to break the fragile glass ampoule of SOz. The relative humidity in the resulting SO,-air-water vapor mixture at - 22°C was - 90 ‘i,; the SO, concentration molar ratio was - 31. The was - 650ppm; the H,O/SO, time allowed for isotopic equilibration ranged from 0.5 to
2 - mm
PLATINUM WIRE’ ELECTRODES
Fig. 2. Glass reaction chamber. 20 h, with results indicating that c 0.5 h was adequate to achieve essentially complete equilibration. Two procedures were used for separating the SO, from the air and water vapor in the equilibrated mixtures. In one procedure, the air mixture was exhausted from the chamber through the following traps: a dry-ice cold trap ( - 79°C) for removal of practically all of the water vapor, a trap containing Mg(CIO,), for removal of the last traces of water vapor, and a liquid-nitrogen cold trap (- 196°C) for removal of the SO,.
Oxygen-18 study of nonaqueous-phase The SO2 was quantitatively recovered and stored in a sample bulb for subsequent isotopic analysis (Holt, 1977). The other procedure was the same except that all of the water vapor was removed at room temperature (* 22°C) by exhausting the air mixture through Aquasorb (No. 6063, Mallinckrodt, Inc.), a chemical desiccant (P,O,) mixed with an inert substance. Oxidation of SO2 induced by a high-voltage spark
Oxidation of SOz in humidified air was induced by a spark discharge generated by a Tessla-coil high frequency generator (Model BDlO, Electra Technic Products, Chicago, IL) in the 3.1-P glass chamber (Fig. 2). The two platinum electrodes provided a spark gap of - 1 cm. The two stopcocks provided a pipet for sampling the gas mixture so that subsequent determinations of SO, concentration could be made by mass spectrometry. Fifty microliters of water was sealed in a fragile glass capsule and placed in the 3.1-t chamber. The chamber was evacuated, 9O~niol of SO2 were cryogenically added, and the chamber was filled to atmospheric pressure with CO,-free, moisture-free air. The ampoule was broken to release the water vapor, yielding a humidity of 75595x, and the gas mixture was allowed to stand until isotopic equilibration between the SOa and the water vapor was attained (at least 0.5 h). The Tessla coil was activated to generate a spark across the electrodes for varying lengths of time. A lO-min spark was sufficient to induce complete oxidation during a subsequent 16-h (overnight) period. Deposits of very finely divided liquid particles appeared on the inner walls of the chamber a few minutes after the relatively brief exposure to the electric spark. After - 16 h, the oxidized product was rinsed from the reaction chamber into a beaker for subsequent isotopic analysis (Holt, 1977). Oxidation by NO2 addition Ninety micromoles of tank NO, were sealed in a glass ampoule and placed in the reaction chamber; 90pmol of SOI were measured and stored in a U-tube trap connected to a vacuum line via a four-way stopcock; and 50$ of liquid water of known 6isO were cryogenically transferred to the reaction chamber and evaporated. The chamber was filled with CO,-free, water-free air, which was transported to the chamber through the U-tube trap in which the SO2 had been stored. Reaction of the NO* with the SO,-air-H,O(g) mixture was started by shattering the NO, ampoule. In some experiments, the subsequent decrease in SOr concentration in the reaction chamber was monitored by periodic mass spectrometric analyses of the air mixture. The tank NO, was not assayed, it was separated from NO with which it may have been associated in the tank by cryogenic transfer to the glass ampoule. Oxidation with gamma irradiation For y-ray irradiation, the high-voltage spark procedure was followed except that, instead of operating the Tessla-coil spark, the 3.1-e reaction chamber was placed in an irradiation cave for a lOO&min (overnight) exposure to y-rays (- 4 kradmin-‘) emitted from a 6oCo source (Blomgren, 1958). Oxidation on charcoal In some experiments, SO, was coadsorbed with water vapor and air on activated charcoal without prior isotopic equilibration between the SO, and water vapor; in others, pre-equilibrated mixtures of Sot, air and water vapor were exposed to dry charcoal for coadsorption. Activated coconut charcoal (8-12 mesh, CX640-L996, supplied by Matheson, Coleman and Bell, Norwood, OH) was used in both procedures. Since pre-equilibration affected the oxygen isotopy of the sulfate formation, the two procedures are described separately.
oxidation of sulfur dioxide
SO2 unequilibrated with water vapor. A 3-g bed ofcharcoal was placed in a closed, 1-e loop of recirculated humidified air and maintained at 110°C to ensure absence of liquid water. By continuous recirculation of the air at atmospheric pressure through a bubbler of water of known 6’*0, the water that adsorbed on the heated charcoal was allowed to reach isotopic equilibrium with the water vapor in the air. Sulfur dioxide of known 6i80 was then loaded onto the charcoal bed at llO”C, without prior exposure to the water vapor. Recirculation was then continued for 15 min, after which the charcoal was emptied into a beaker and covered with 50 ml of leach water of known 6180. One milliliter of 1: 1 HCI was added and the solution was heated to boiling and filtered through Whatman No. 42 filter paper, retaining the charcoal in the beaker. These steps were twice repeated, using 15-20 ml of water and 0.5 ml of 1: 1 HCI. The filtrates were combined and reserved for subsequent isotopic analysis. Pre-equilibrated SO,-air-water vapor. For one set of samples, the 3-g bed of charcoal was preconditioned by vacuum heating at 300°C followed by equilibration at 110°C with air and with water vapor of the same 6i80 to be used in the experiment. A mixture of SO,, air and water vapor, preequilibrated in the reaction chamber, was then evacuated from the chamber through the 3-g bed of charcoal at 1 lO”C, a cold trap at - 79°C and a cold trap at - 196°C. The SO, was collected on the charcoal and the water vapor was collected in the trap at -79°C. No SO1 was detected in the trap at - 196°C indicating quantitative retention by the charcoal. For another set of samples, the charcoal was preconditioned only by vacuum heating at 300°C. The SO,-air-water vapor mixture was then passed through the charcoal bed at room temperature (- 22°C). Under these conditions, the charcoal was not pre-equilibrated with the water vapor; and, at room temperature, some of the water vapor was retained on the charcoal as well as in the dry-ice cold trap.
3. RESULTS AND DISCUSSION
Isotopic equilibration of SO2 and water vapor In Fig. 3, the 6r80 of SO2 is plotted vs the 6 ls 0 of the water vapor. Curve A represents the experiments in which the water vapor was cryogenically separated at - 79°C from the SO,air-H, 0 (g) mixture; for curve B, the water vapor was chemically separated at 22°C. Regression analyses of the data yield the following equations: curve A, 6is0 so1 = 1.01 6i80
curve B, 6r*O so2 = 0.94 Lii*o
H20(g) + 24.2
The average 6” 0 for five samples of the original SO2 was 13.7 f 1.2 %,. The 8i8 0 values obtained from two control runs (same procedure but with dry air) also averaged 13.7 %, These results strongly indicate that rapid isotopic equilibration occurred between SO, and water vapor under the experimental conditions. In the equilibration reaction, SOz (g) + Hz O(g) *SO, . Hz 0, the structural form of SO,. H,O (e.g. gas-phase or surface-adsorbed) is unspecified. The slope of approximately 1 in (1) and (2) indicates that the 6 r * 0 of the oxygen atoms in the equilibrated SO, was completely dependent upon the S l8 0 of the water vapor to
B. D. HOLT et al.
terminated, and that NOz was probably a major reactant in the oxidation of SOZ to H,S04, presumably through a chain reaction mechanism. (The species O,, HzOz, OH or HO, were not detectable by the mass spectrometer used.) The depletion of SO, in each experiment was accompanied by a corresponding appearance of very finely divided liquid droplets on the inner walls of the chamber. Evidently these were droplets of aqueous H,SO,. Unlike water, the droplets were not easily evaporated by increasing the temperature a few degrees; and nearly quantitative recovery of the sulfur, introduced as SO,, was obtained as sulfate in the solution obtained by rinsing the chamber. The isotopic data obtained in these experiments fit the relationship: B1*Oso:~ = 0.69 #*OHZO(gI+ 15.7%0.
10 I A. Cryogenic B: Chemical 0 . . ..I..... -10 -20
Separation at -79’C Separation at 22’C ..,I.... 0 10
Fig. 3. SO,-water
it was exposed. The difference of - 19 ‘i,, between the y-intercepts (Fig. 3) is in approximate agreement with the difference in the thermodynamic fractionation factors that are calculable from spectroscopic data for the gaseous species, SO, and H, O(g), at the two temperatures - -79°C and 22°C (unpublished data). This agreement was evidently caused by rapid isotopic exchange between SO, and H,O(g) within the inlet zones of the dry-ice cold trap (- 79°C) through which the gas mixture was drawn in the first set of experiments (curve A). These results have the following implications: (1) SO, in the atmosphere rapidly equilibrates with ambient water vapor; as a result, its 6180 is dynamically controlled by the al80 of ambient water vapor, regardless of the BL80 of the SOZ at its point of origin. (2) Measurement of the 6”O of SO, cannot, therefore, be used to determine its source of emission. (3) The 618 0 of secondary sulfates, formed in the atmosphere from SO, and water vapor, does not depend upon the al80 of the SO, as an independent variable. which
of SO, in air with high-volrage
Kinetic data were obtained by mass spectrometric monitoring of the temporal depletion of SO, in the SO,-air-water vapor mixture, after a relatively brief exposure to the high-voltage Tessla-coil spark. These mass spectrometric analyses, not reported in this paper, indicated that the spark generated one or more gaseous compounds which caused oxidation of the SO, to continue for several hours after the spark was
For the purposes of representation in Fig. 1, all experimental data discussed in this paper were converted to the liquid water basis, using the following 6’*0 relationship between water vapor and precipitating liquid water at 20°C (Broecker and Oversby, 1971): 8180H20(1, = @80”,o(p)
Curve 4 in Fig. 1 represents the result of the substitution of (4) into (3). There is also another small correction that arises from the finite amount of water vapor available for the reaction. This results in a change in the a’*0 of water vapor as the reaction proceeds. For our experimental conditions, where [HZO]/[S02] - 31, this correction amounts to - 6 “/;I. When increased by this 6 ‘>”correction, the slope of curve 4 approximates a ratio of 3/4, indicating that the isotopy of three of the four oxygen atoms in the sulfate was controlled by the 6’*0 of the water vapor. A mechanism of transformation of SO, to sulfate satisfied by this condition is isotopic equilibration between the SO2 and excess water vapor, followed by oxidation that incorporates one additional oxygen from the water vapor and one oxygen from the oxidant, so that three of the four sulfate oxygens are controlled by the water vapor. A 3/4 control of sulfate oxygens by water is also characteristic of aqueous-air oxidation mechanisms in which the HSO; ion isotopically equilibrates with liquid water before oxidation to SOi- (Holt et al., 1981a). However, the mechanisms responsible for the two transformation processes necessarily are entirely different. Oxidation with NO* When NO, was added to mixtures of SO, and humidified air, the rate of depletion of SO, and the liquid condensate formed on the chamber walls were very similar to those of the high-voltage spark experiments. Curve 5 in Fig. 1 shows the S180s02_ vs 6’80 H,O(I)relationship derived by substitution Gf (4) into the water-vapor relationship: 6180so:~
= 0.84 618OfilO(gI+ 13.8 “/o,.
Oxygen-18 study of nonaqueous-phase
Curve 6 in Fig. 1 was derived by substitution into the water-vapor relationship:
The higher slope of (5) than (3), 0.84 vs 0.69, suggests a greater dependence of the 6180 of sulfate on the 6180 of the water vapor when the NO2 is added directly rather than being produced by an electric spark. This greater dependence is probably related to more isotopic exchange between NOz and H,O(g), prior to complete oxidation of the SO,. Surfacecatalyzed reactions of NOz and Hz0 to produce HNO, and HNO, are well known. Oxidation
6180 so:- = 0.71 8’~OH10,g)f Oxidation
sl*o so42- = 0.2286’80
10 6 6 soz 4 ___:
PO -29 15
Water Vapor (“/O”) -14
. . ..~..~...1~,.,~./.,..,....,~.. --15
with water vapor. Figure 4 shows results for the oxidation of unequilibrated SOz on dry charcoal. Curve “SO,” represents 6180sojm vs 6l80 so2 at constant #*OHlo,g, (- 16.9 y,“,) and 6l*O leach wa,er (-7.9 Ti,); curve “water vapor” is POSO’~ vs #go H,O(gjat constant #*Oso, (10.5 %,) and 6 lableachwater(- 7.9 %,); and curve “leach water” is 6’80 *- vs PO leach water at constant ‘8oS0, (1o.5 %,) and h’30 H,O(gj(- 16.9 %,). The equations that fit the experimental data are:
in air with gamma irradiation
In mixtures of SO1, air and water vapor, the SO, concentration decreased - 8 “/;Iduring 60 min of exposure to y-irradiation; after 1000 min of exposure, it decreased to > 95 7,. The lOOO-min exposures were used in four experimental runs. The inner walls of the glass chamber were dry before irradiation and coated with a deposit of finely divided liquid droplets (H, SO, solution) after irradiation.
oxidation of sulfur dioxide
Fig. 4. Oxidation of SO2 on charcoal at 110°C. SO, was not isotopically pre-equilibrated with water vapor.
- 4.4 o,,,,
B. D. HOLT et al.
for SO*, 6isO soi- = 0.441 61sOH20(g) + 5.4%”
for water vapor, and 6180so:-
= - 1.8 %,
for leach water. The zero slope for the leach water curve indicates that sulfur(IV completely oxidized to sulfur(V1) on the charcoal surface (at 1lOC) prior to contact with the leach water. One implication of this observed dry oxidation on charcoal, as it relates to the formation of atmospheric sulfates, is the importance of heterogeneous oxidation of SO, to sulfate on active surfaces of carbonaceous particles prior to the particles’ subsequent association with liquid water (clouds, fog, rain, etc.). To the extent of occurrence of such a mechanism, involvement of aqueous-phase, carbon-catalyzed oxidation mechanisms is necessarily precluded. The slopes of - 2/9 (0.228) and - 4/9 (0.441) for the SO, and water-vapor curves, respectively, support the hypothesis that the SO, combines and isotopically equilibrates with air oxygen and water vapor in an intermediate molecule. This molecule is chemisorbed on the surface of the charcoal according to the relationship, 6i80so:~
where (for our experimental conditions) C = 3.1 yoo,. The intermediate molecule apparently consists of nine oxygen atoms, two of which are isotopically controlled by the SOZ and four are controlled by the water vapor. The 9-oxygen intermediate molecule decomposes in leach water, yielding sulfate ions of unchanged isotopic enrichment. A structural form of the chemisorbed molecule consistent with these observations is
SO2 pre-equilibrated with water vapor. Figure 5 shows the dependence of the 6180 of sulfate on the 6” 0 of water vapor when the SOZ is pre-equilibrated with the water vapor in air before it is adsorbed and oxidized on dry charcoal. Curve A represents a calculated relationship for charcoal at llO”C, obtained by substituting (2), with the approximated slope of 1.0 and the y-intercept adjusted for llo”C, into (10). This curve is given by: 8’80so:-
Curve B represents the experiments in which the charcoal was preconditioned with water vapor at 110°C and was then held at 110°C during subsequent adsorption of SO2 from the SOZPairPH, O(g) mixture. This curve is given by: 6i80so:~
= 0.61 6180H,0,g,+3.8’;,0.
8180s042- = 0.536180H20,g,+10.8(;oo.
The deviation of the measured Curve B from the calculated Curve A is primarily caused by the small isotopic change in the 618 0 of the water in the closed system of ratio H,O/SO, = - 31. The isotopic shift of - l’:“c between Curves B and C was probably caused by the difference in temperatures, as well as by
A Calculated. B: Measured, C: Measured.
1lO’C 1lO’C ZX
Curve C represents the experiments in which the charcoal was not preconditioned with the water vapor and was held at room temperature (- 22°C) during adsorption of the SO, from the SO,-air-H,O(g) mixture. This curve is given by:
Curve Curve Curve
Fig. 5. Oxidation of pm-equilibrated SO,-air-water
vapor on charcoal
Oxygen-18 study of nonaqueous-phase the different role of water on the surface of the charcoal at the two temperatures. Liquid water was undoubtedly present in the pores of the charcoal surface in the experiments of Curve C, but not in the experiments of Curve B (Smith, 1959). Curve C, which more closely simulates atmospheric conditions, was combined with (4) and plotted as Curve 7 in Fig. 1.
4. COMPARISON OF ISOTOPIES OF SULFATE FORMATION Figure 1 shows the 6180so:m vs 6’80H,0 relationship obtained from seven methods of sulfate preparation. Broken lines (curves l-3) represent aqueous preparations previously reported (Holt er al., 1981a); solid lines (curves 4-7) are for the non-liquid preparations of this report. Also plotted in Fig. 1 are the isotopic data obtained for samples of precipitation water (rain or snow) collected at Argonne, IL. during the neriod of October 1976 to March 1978 (Halt et-al., I981b). All of the precipitation samples contained sulfates that were higher in S”O than the sulfates formed by the seven laboratory-tested mechanisms. It appears, therefore, that the precipitation sulfates did not originate solely via any of these mechanisms, either singly or in combination. Apparently, one or more components of “0 sulfate were present in the highly enriched atmosphere which when mixed with sulfates of lower ‘*O content, formed by one or more of the mechanisms listed above, gave the observed 6180 values in precipitation sulfates. Current studies indicate that primary sulfates are generally much higher in 6”O than secondarv sulfates and that their relative amounts in the atmosphere can be estimated by isotopic analysis. Such estimates have recently been reported (Holt et al., 1982).
oxidation of sulfur dioxide
reactions pertaining to HISO, aerosol formation. Int. J. Chem. K&et., Symp. 1, 629-640.
Castleman A. W., Jr. and Tang I. N. (1977) Kinetics of the association reaction of sulfur dioxide with the hydroxyl radical. J. Photochem. 6, 349-354. Clarke A. G. and Britton L. G. (1977) Heterogeneous reactions of sulfur dioxide and nitrogen oxides on carbon. Proc. Fourth Int. Clean Air Congr., Japanese Union of Air Pollution Prevention Associations. Corn M. and Cheng R. T. (1972) Interactions of sulfur dioxide with insoluble suspended particulate matter. J. Air. Pollur. Control Ass. 22, 870-875. Cox R. A. (1973) Some experimental observations of aerosol formation in the photo-oxidation of sulfur dioxide. Aerosol sci. 4, 473483. Davis D. D. and Klauber G. (1975) Atmospheric gas phase oxidation mechanisms for the molecule SOZ. Int. J. Chem. Kinetics, Symp. 1, 543-556. Davtyan 0. K. and Ovchinnikova E. N. (1955) Chemisorption and the oxidation of sulfur dioxide on solid catalysts at room temperatures. Doklady Akad. Nauk S.S.S.R. 104, 857-860, through CA 50, 9844b (1956). Dietz R. N. and Garber R. W. (1978) Power plant flue gas and plume sampling studies. Progress Report No. 2 August 1977 to Julv 1978. BNL No. 25420, Brookhaven National Laboratory, Upton, NY. Dietz R. N. (1979) Factors affecting the concentration of H,SO,/MSO, in combustion gas. Presented at Engineering Foundation Conference on Stack Sampling and Stationary Source Emission Evaluation, 1 April 1979, unpublished work, Brookhaven National Laboratory, Upton, NY. Evans F. D. and Targett B. H. (1978) Some applications of the ultramicrobalance to power station flue gas research. Thermochim. Acta 1978, 281-291.
Gordon S. and Mulac W. A. (1975) Reaction of the 0H(x2rr) radical produced by the pulse radiolysis of water vapor. Int. J. Chem. Kinetics, Symp. 1, 289-299. Haury G., Jordan S. and Hofmann C. (1978) Experimental investigation of the aerosol-catalyzed oxidation of SO, under atmosoheric conditions. Atmosoheric Environment 12, 281-287: Holt B. D. (1977) Preparation of carbon dioxide from sulfates, sulfur dioxide, air, and water for determination of oxygen isotope ratio. Analyt. Chem. 49, 16641667. Holt B. D., Kumar R. and Cunningham P. T. (1981a) Oxygen18 study of the aqueous-phase oxidation of sulfur dioxide. Atmospheric Emironment 15, 557-566.
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(soot) particles. Science, Wash. 186, 259-261. Novakov T. (1978) Role of carbon soot in sulfate formation. Lawrence Berkeley Laboratory, Berkeley, CA, Report No. LBL-7887. Pierce C., Smith R. N., Wiley J. W. and Cordes H. (1951) Adsorption of water by carbon. J. Am. them. Sot. 73, 45514557. Rosen H. and Novakov T. (1978) Identification of primary particulate carbon and sulfate species by raman spectroscopy. Atmospheric Environment 12, 923-927. Schroeder W. H. and Urone P. (1978) Isolation and identification of nitrosonium hydrogen sulfate (NOHS04) as a photochemical reaction product in air containing sulfur dioxide and nitrogen dioxide. Enuir. Sci. Technol. 12, 545-550. Schryer D. R., Coefer W. R., III and Rogowski R. S. (1979) Correspondence, Enoir. Sci. Technol. 13, 1419.
Shen C. H. and Springer G. S. (1976) Photochemical particulate formation in sulfur dioxide-air mixtures. Atmospheric Environment 10, 235-239. Smith R. N. (1959) The chemistry of carbon-oxygen surface compounds. Q. Rev. 13, 2877305. Stopperka K., Wolf F. and Suess G. (1968) Gas-phase reaction of nitrogen dioxide and sulfur dioxide. Z. anorg. al/g. Chem. 1968, 1429. Tartarelli R., Davini P., Morelli F. and Corsi P. (1978) Interactions between SO2 and carbonaceous particulates. Atmospheric Environment 12, 289-293.
Toossi R. (1978) Physical and chemical properties of combustion generated soot. California University, Berkeley, CA. Lawrence Berkeley Laboratory Report LBL-7820. Wood W. P., Castleman A. W., Jr. and Tang I. N. (1975) Mechanisms of aerosol formation from SOz. J. Aerosol Sci. 6, 3677374.