Environ. Sci. Technoi. l W 3 , 27, 2162-2167
Oxidation of Aqueous Sulfite Ion by Nitrogen Dioxide' David Littiejohn, Ylzhong Wang,* and Shih-Ger Chang'
Lawrence Berkeley Laboratory, Energy and Environment Division, Berkeley, California 94720 The reaction of nitrogen dioxide with aqueous sulfite solutions, in the presence and in the absence of oxygen, has been studied using Raman spectroscopy to determine product concentrations. The products observed include nitrite, sulfate, and dithionate ions. The reaction appears to initially produce nitrite ion and sulfite radical: NO2 + S032- NO2- + SO3*-. The sulfite radical can undergo either recombination or reaction with oxygen to form SO6*-. In the absence of oxygen, we obtain a ratio of [so42-]/[s2062-] = 1.8 f 0.3, reflecting the branching of the recombination of sulfite radical: SO3*- + SOB*S 2 0 6 2 - or SO3*- SO3*so32-+ SO3. The production of SO3'- and SO6'- radicals from the nitrogen dioxidesulfite ion reaction suggests that this reaction could contribute significantly to S(IV) oxidation in atmospheric aerosols under suitable conditions.
-
-
-
+
Introduction Reactions involving nitrogen oxides and sulfur dioxide are of considerable importance both in air pollution and in control of emissions from coal-fired power plants. Nitrogen dioxide and nitric oxide are both free radicals and are quite reactive. Nitrogen dioxide is much more soluble than nitric oxide, and it can react with water to form nitrite and nitrate ions:
+
+
N02(g) N02(g) H 2 0
-
2H'
+ NO; + NO,
(1)
NO2 can also react with a number of species in aqueous solution. One reactant of particular interest is dissolved sulfur dioxide, especially when it is in the form of either hydrogen sulfite ion or sulfite ion. There have been a number of studies of the reaction of nitrogen dioxide with sulfur dioxide dissolved in aqueous solution (1-9). However,the agreement between the studies is not always good, and some of the details of the reaction have not been established. Investigation of the reaction has been hampered by the reaction of nitrogen dioxide with water (10)) the dimerization of NO2 (11))and the presence of impurities. There is disagreement about the effect of NO2 on the autoxidation of sulfite. This is due, in part, to the lack of information on the products of the aqueous reaction of nitrogen dioxide with sulfite and hydrogen sulfite ions. It has been proposed (9) that the reaction proceeds via one or more intermediates to form nitrite and sulfate ions. The overall reaction appears to be (9) 2N02 + SO-: 2N0,
+ H20
-
+ HSO, + H 2 0
2H'
+ 2NO; + 50;-
(2)
3H'
+ 2NO; + SO-;
(3)
-
The investigation of the reaction under acidic conditions
* Author to whom correspondence should be addressed. t This work was supported by the Assistant Secretary for Fossil Energy, Office of Coal Utilization Systems, U.S.Department of Energy, under Contract No. DE-AC03-76SF00098 through the Pittsburgh Energy Technology Center, Pittsburgh, PA. t On leave from the Research Center for Environmental Sciences, Acadernica Sinica, Beijing, Peoples Republic of China.
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Envlron. Scl. Technol., Vol. 27, No. 10, 1993
is complicated by the reaction of nitrite ion with hydrogen sulfite ion to form a number of nitrogen sulfonates (12). We have investigated the products formed by the reaction of nitrogen dioxide with aqueous sulfite ion, both in the presence and in the absence of molecular oxygen. From the products observed, we propose a reaction mechanism that is consistent with recent kinetic studies.
Experimental Section The reaction was investigated by flowing gas mixtures containing nitrogen dioxide over aqueous sulfite ion solutions in a small long-necked bulb. Small samples of the solutions were collected periodicallyand analyzed using Raman spectroscopy. The gas mixture flowed through an inner tube that was coaxial with the neck of the bulb, and then over the surface of the solution and out of the neck. Two similar cells were used. The solutions were stirred at approximately 200 rpm, which provided mixing without creating significant turbulence at the solution surface. The larger cell was used with 10 mL of solution and a gas flow rate of 1.4 L min-'. The solution in this cell had a surface area of approximately 9 cm2. The smaller cell was used with 5 mL of solution, which had a surface area of about 6 cm2. Gas flow rates of either 0.5 or 1L min-1 were used with the smaller cell. With the conditions used, generally less than 20% of the NO2 that flowed into the cell was absorbed by the solution. However, at the lowest nitrogen dioxide concentrations used, absorption reached 25%.The inlet and outlet concentrations of nitrogen dioxide in the gas flow could be monitored by a Thermoelectron Model 14A chemiluminescent NO, analyzer. The solutions were prepared with house deionized water that had been passed through another deionizing column and stored in aplastic container. The cells were washed with acid prior to use and rinsed with deionized water. EDTA (1mM) was added to most solutions to minimize interference from metal ions. Reagent grade chemicalswere used in solution preparation. The gas mixtures used in this study were prepared by diluting mixtures of NO2 in N2 with research grade nitrogen and oxygen when desired. Mixtures of 914 ppm and 2.6 % NO2 in nitrogen (Matheson) were using in preparing the gas mixtures for the experiments. The reaction was studied using oxygen concentrations ranging from 0 to 5 % . Samples of the solutions under study were collected periodically in 1-mm-diameter glass tubes for analysis by Raman spectroscopy. Raman spectra of the samples were obtained using a Spex 1403 double monochromator and a Spex Datamate data aquisition system. The samples were illuminated with about 100 mW of 514-nm light from a Coherent 90-5 argon ion laser. The concentrations of the anions of interest were determined by comparison of their peak areas with those of an internal reference added to the solution (13). Reference ions used included perchorate ion, boric acid, nitrate ion, and carbonate ion. Some runs were done without references to insure that the reference compound did not significantly interfere with the reaction. Most of the measurements were made at 20 OC, although a few 0013-936X/93/0927-2162$04.00/0
0 1993 Amerlcan Chemlcal Soclety
agreed quite well with the loss of sulfite ion, indicating a good mass balance for sulfur. The data from our experiments are compiled in Table I. Dithionate ion was observed as a product in all of the reaction mixtures from studies done without oxygen. Sato et al. ( 4 ) had also observed dithionate in the product mixture from the reaction of sulfite ion with nitrogen dioxide, Most of the previous studies did not identify the products from the nitrogen dioxide-sulfite ion reaction. A feasible mechanism for the production of dithionate is the recombination reaction of sulfite radical (14):
-
so,'- + so,'-
I
1
750
(44
A second branch of this reaction also occurs (15):
,
950
S206"
1150
+ SO,'SO, + H 2 0
SO,'-
Raman shift (cm-')
Flgure 1. Raman spectrum of a solution of 1.0 M SOS2-and 0.030 M C10,- exposed to 5000 ppm NOp In Np for 1 h.
runs were conducted at 38 "C and 50 "C. Experiments were done with gaseous nitrogen dioxide concentrations ranging from 0.1 to 9000 rpm. The sulfite ion concentrations used ranged from 0.10 to 1.0 M. The high sulfite ion concentrations were required to obtain reasonably strong Raman signals and to minimize interference from the competing reaction of nitrogen dioxide with water. The solutions studied ranged from mildly acidic (pH 4) to highly alkaline (0.8 M NaOH). In the acidic solution, investigation of the reaction was complicated by the reaction of hydrogen sulfite ion with nitrite ion to produce nitrogen sulfonate compounds. The concentrations derived from the Raman spectra were used to derive profiles of concentrations with time of the species of interest. A few of the reaction mixtures were analyzed for nitrite ion by ion chromatography after exposure to NO2. No attempt was made to determine the rate constant for the reaction of nitrogen dioxide with sulfite ion, since reactant diffusion and mass transfer may limit the kinetics in the system used, and the kinetics have been investigated by Clifton et al. (9) and by Huie and Neta (8). Clifton et al. reported a rate constant of 1.24 X 107 M-ls-l at pH 5.3 and 2.95 X 107 M-l s-l at pH 13.
-
Results and Discussion Measurements in the Absence of Oxygen. The reaction of nitrogen dioxide with sulfite ion was initially studied in the absence of oxygen. A Raman spectrum of a solution of 1.0 M sulfite ion and 0.030 M perchlorate ion exposed to 5000 ppm NO2 in nitrogen flowing at 1.4 L/min for 1 h is shown in Figure 1. Perchlorate ion was added to the solution to act as a reference for the determination of concentrations from the Raman spectrum. The nitrite ion peak is weak because nitrite ion does not Raman scatter very strongly compared to the other species. The peak areas of the species of interest were corrected for their scattering intensities and used to derive the concentrations of each species. There is some uncertainty in the determination of low concentrations of the products, particularly with species such as nitrite ion that do not scatter strongly. Therefore, slopes of plots of concentration as a function of time were found to be a better measure of the reaction progress than the individual concentration measurements. Values for the slopes were obtained by linear least-squares fib. The buildup of sulfate ion and dithionate ion generally
-
-
SO:-
2H'
+ SO,
+ SO-:
(4b) (5)
The branching of this reaction has been studied by Eriksen (16) and, more recently, by Waygood and McElroy (17). Eriksen observed a pH dependence in the production of sulfate and dithionate ions. Waygood and McElroy used a less complex chemical system and derived a value for the branching ratio of k 4 J k 4 b = 0.8 f 0.2 at pH 4.3. The sulfite radical could conceivably be formed by the reaction of sulfite ion with nitrogen dioxide:
NO,
+ SO,"
-
+ SO,'-
-
+ SO,'-
(6) In the absence of other significant reactions producing sulfate ion or dithionate ion, the ratio of [S04"1/[S2062-] obtained from our measurements should be similar to the value determined by Waygood and McElroy (17). A plot of [ s o 4 2 - ] / [ s 2 0 6 2 - ] as a function of solution pH is shown in Figure 2. The values of our measurements from the reaction of NO2 with sulfite ion in the absence of oxygen using gas-phase NO2 concentrations of 1000-5000 ppm and SO3,- concentrations of 0.1-1 M are shown in Figure 2, along with those derived from Eriksen's work (16) and from Waygood and McElroy's measurements. It would not be expected that a radical recombination reaction such as reaction 4 would be substantially influenced by the pH of the solution. I t appears that excess sulfate ion was produced in Eriksen's experiments by other reactions, particularly at the less acidic conditions. The [ S 0 4 2 - ] / [S2Q62-] ratio from our measurements agrees reasonably well with that value given by Waygood and McElroy, although they obtained a ratio that is lower than the average of our measurements. A reaction such as 7,
NO,
NO,
NO;
+ SO,
(7)
combined with reaction 5, could account for the excess sulfate that we observe. The appearance of nitrogen sulfonate compounds in the reaction mixture at pH 5 and below prevented us from obtaining a measurement at the conditions used by Waygood and McElroy. The average value of [ S 0 4 2 - l / [ S 2 0 ~ 2 - ] from our measurements is 1.8 f 0.3. The rate of production of the reaction products was measured as a function of the gaseous nitrogen dioxide concentration using solutions of 1M sulfite ion and 0.8 M sodium hydroxide. The [ S 0 4 2 - 1 / [ s 2 0 6 2 - ] and [N02-1/ [ S 0 4 % 1 product ratios were calculated using their respective rates of production. These ratios are plotted as a function of gaseous nitrogen dioxide concentration in Environ. Sci. Technol., Vol. 27, No. 10, 1993 2163
Table I. Results of Measurements init [S(IV)I (M)
[Nod, gas (ppm)
PH
d[S03'-l/dt (M 9-l)
d[SOPl/dt (M 8-l)
d[Sz0s21/dt (M d) d[NOz-lldt (M8-1)
No Oxygen 5 1400 1400 1400 5 1000 3000 5000 7000 9000
0 0.11 0.11 0.15 0.26 1.0 1.0 1.0 1.0 1.0
9.5 9.0 8.0 5.5 9.7 13.9 13.9 13.9 13.9 13.9
0.25 1.0
5 5000
9.5 13.9
1.0
5000
13.9
0.20 0.25 0.25 0.25 0.25 0.25 0.26 0.26 0.26 0.26 0.26 0.26 0.26 0.26 0.26 0.40 0.60 1.0 1.0 1.0 1.0
5 5 2.0 2.0 2.0 5 0.1 0.3 0.3 0.3 5 10 13 23 50 5000 5000 1000 3000 5000 5000 5000 5000
9.0 9.5 9.0 9.0 12.0 9.0 9.0 9.0 9.0 9.0 13.1 9.0 9.0 9.0 8.5 13.9 13.9 13.9 13.9 13.9 13.9 13.9 13.9
1.0
5000
13.9
1.0 1.0
l4
1.8 X 10-6 7.0 X lo4 1.8 X 10" 1.2 x 10-5 -6.3 X 10-6 -2.3 x 105 -3.3 x 10-5 -4.9 x 105 -6.4 X 10"
2.6 X 10.0 x 1.4 X 2.2 x 3.3 x
4.6 X 10-6 8.9 X 10-6 4.9 x 10-6
10-6 10-6 10-5 105 10-5
5.6 X 8.9 X 1.3 X 1.6 X
1%Oxygen -6.1 x 10-6 5.4 x 10-6 -5.6 X lo4 3.4 x 105 2% Oxygen
10-5 4.6 X 106 3% Oxygen -1.6 X 10" 1.6 X 10" 1.2 x 10-5 -9.0 x 10-6 1.1x 104 1.3 X 106 -1.5 x 105 -1.1 x 105 9.8 X 10-6 -1.8 X 2.1 x 10-5 -6.9 X 10-6 6.9 X lo4 -5.7 x 10-6 6.0 X lo4 -1.2 x 105 1.1x 10" -1.04 X 1.03 X 106 -2.4 x 10-5 2.4 X 10-5 -1.5 X 10" 1.6 X 106 -2.7 x 105 2.7 X 106 -2.6 X 10" 2.8 X 10" -2.6 x 10-5 2.8 X 10" 3.1 x 10-5 -4.3 x 10" -7.3 x 105 5.7 x 105 -3.7 x 10-5 3.3 x 105 -5.9 x 10-5 4.9 x 10" -5.6 X 106 3.4 x 10" -6.6 X 10-5 4.6 X 10-5 -8.1 x 10-5 6.4 x 105 -10.7 X 10-5 8.9 x 105 5% Oxygen 8.9 x 10-5 -1.1 x 10-4 -6.6
1.7 x 7.1 X 2.8 X 3.8 X 5.2 X 7.3 x
1.8 X lo4
X
10-6 10-6 1P
106
10-7
10-6 106 106 106 106
1.6 x 10-7 1.1x 105
1.7 x 10-7 6.1 X 106
1.0 x 1od
5.3 x 10-5 1.7 x 10-7
-1.4
X
10-6
9.7 x 10-7
5.2 X 7.5 x 1.9 x 6.0 X 10.6 X 10.1 x 8.5 X 8.4 X
2.0 x 3.4 x 1.1 x 3.9 x 6.1 X 5.3 x 5.1 X 5.1 X
10-6 10-6
10-6 lo4 10-6 1o-s
10-6 10"
8.4 X lo4
10-6 106 106 106 106 106
106 106
5.1 X 10-5
I 4 1
12
-
10
-
W
A
8-
1
6W 4W
W
X
X
A 2
4
6
8
X
10
12
14
PH Flgure 2. Comparison of the [S042-]/[S20e2-]ratios obtained by Eriksen (ref 16, W) and Waygood and McElroy (ref 17, A) from sulfite radical recombination and this study (X) from the nitrogen dioxidesulfite ion reaction.
Figure 3. The [N02-]/ [ SO42-I values are about what would be expected from the reaction of nitrogen dioxide with sulfite ion proceeding via reactions 6,4a, and 4b, using the branching ratio obtained by Waygood and McElroy (17). 2164
Envlron. Sci. Technol., Vol. 27, No. 10, 1993
0
2000
4000
6000
8000
10000
[NO21 ppm
Figure 3. Measuredratios of [S042-]/ [S20e2-](W) and [NO2-] / [ S042-] (A) as a function of gaseous nitrogen dioxide concentration (1.0 M S032-, 0.8 M OH-, no oxygen).
Measurements in the Presence of Oxygen. While measurements were made with gas-phase oxygen concentrations ranging from 1 to 5 % , most of the experiments with oxygen were done using 3 % oxygen in the gas mixture.
Table 11. Redox Couples Related to the Proposed Reaction Mechanism
i
redox couple
A LD v
so3’-/so32so3/so3’so4’-/so42.
A A
a,
,-
a:
NOdN02-
A A A
0 r
2-
ref
1.04
22 24 22 22 22
0.72
0.25 2.43 1.4
HSOs’/HSOg
C
0 .-c
Eo (V)
A A
E
A
6 .
0 .1
..1..1.,
I
I
.I
. I . . ,
I . , . 11.1
I
100
10
1
. ..
I....,
1000
.
8 . -
0 10000
PO21 ppm
Flgure 4. Sulfate ion production rate as a function of gaseous nitrogen dioxlde concentration in the presence of 3 % oxygen (0.20-0.26 M S0s2-, pH 9).
The same cells used in the oxygen-free experiments were used in the experiments with oxygen. The product distribution obtained with oxygen in the gas stream differed significantly from the measurements without oxygen. The proportion of sulfate ion increased relative to both dithionate and nitrite ions. With a gas stream containing 3 % oxygen,the fraction of sulfite ion converted to dithionate ion decreased as the gaseous nitrogen dioxide concentration was lowered. Our detection sensitivity for dithionate was not sufficient to observe dithionate in all of the runs. With nitrogen dioxide concentrations below about 5 ppm, dithionate ion was undetectable. This can be attributed to sulfite radical reacting by a pathway other than reaction 4, presumably due to the reaction with oxygen: SO,’-
+ 0,
-
SO,‘-
(8)
This reaction is very fast, approaching the diffusioncontrolled limit (18). At low sulfite radical concentrations and large dissolved oxygen concentrations, reaction 4 is insignificant compared to reaction 8. Under these conditions, sulfate ion can be produced by other reactions, including 9 and 10:
so,‘- + 50;-
- sot-+ so,’-
+ S032-- SO42- + SO,*SO,*- + SO-: SO:- + SO,*SO;- + H+* HSO; SO-: + 50:HSO; + SO-: SO$-
@a)
-
(9b)
-
(11)
(10)
(12) These reactions have been discussed in other studies of sulfite ion oxidation (14,18,19). The competition between reactions 4 and 8 will cause a variation in the S042-/S20,j2ratio, which we observe. The oxidation of sulfite ion was measured with 0.200.26 M sulfite ion solutions at pH 9 and gas streams of 3% oxygen in which the concentration of nitrogen dioxide was varied. A plot of the rate of sulfate ion production as a function of gas-phase nitrogen dioxide concentration is shown in Figure 4. For comparison with the values shown in the graph, runs with 3% oxygen in nitrogen with no
nitrogen dioxide yielded sulfate ion production rates of about 6 X lo-’ M s-l. This rate of oxidation is at least 1 order of magnitude less that the oxidation rates obtained with nitrogen dioxide present. At low nitrogen dioxide concentrations, the oxidation rate shows a dependence on the nitrogen dioxide concentration. Under these conditions, nitrogen dioxide absorbed by the solution can initiate a free-radical oxidation cycle that can oxidize a number of sulfite ions. At high nitrogen dioxide concentrations, the oxidation rate levels off and appears to be limited by transfer of oxygen into solution. The data in Table I indicate that the sulfite ion oxidation rate has a linear response to oxygen concentration when other experimental variables are fixed. From the observed results, we believe that the overall mechanism for the reaction of nitrogen dioxide with sulfite ion is as follows:
-sos- + - sod+ --sop + + so,.- + - + + + - +
NO2 t SOS” X NO*-t S0s’S0s’- + SO3.SZOS” S0s’- t SO.$- SOSh + SO8 SOS + HZ0 2H+ + so4” +
0%
s05- SOS” + S0s’SOs*- SOs% SO4” SO,’SOS” 50,s 503.-
SO,” H+ HSOi HSOB SO8” SO,” f
sod”
k = 2 X lo7 M-l s-1 k.+b = 5.7 X 108 M-1 a-l
ref 9 ref 18
assumed to be rapid k = 1.1 X 1Og M-’8-1 k.+b = 1.3 X 107 M-1 s-l
ref 18 ref 20
k = 2 X 1Og M-1 E-1 pK = 9.4 k is pH dependent
ref 20 ref 19 ref 21
Rate constants from the literature are included where available. The redox couples for some of the species in the mechanism are listed in Table 11. All of the radical species in the mechanism have been observed, and many of the rate constants have been measured. There are undoubtedly other reactions and species in the reaction system, but they are expected to be less important than those listed above. This mechanism is consistent with our observed product distribution and product dependence on reactant concentrations. Due to mass transfer limitations, it is difficult to obtain a quantitative comparison between the experimental measurements and the proposed mechanism. I t is not clear whether the transient intermediate in reaction 6 is similar to that proposed by Clifton et al. (9). One of their motivations for proposing an intermediate in the nitrogen dioxide-sulfite ion reaction mechanism was a study of electron-transfer reactions by Huie and Neta (8). The Cl02 + reaction proceeds primarily by electron transfer (231, and from Huie and Neta’s measurements, the rate constant for an electron-transfer reaction between NO2 and SO,” would be expected to be more than 1 order of magnitude smaller than the rate constant that they obtained. Thus, the reaction does not appear to proceed by electron transfer. However, Sarala et al. (24)investigated electron-transfer reactions between sulfite ion and metal ion complexes. They noted that sulfite ion reacts more quickly with Mn04- than would be expected by an outersphere electron-transfer reaction. They speculated that this reaction could proceed by Environ. Sci. Technd., V d . 27, No. 10, 1993 2185
Table 111. Oxidation of S(IV) by NOa and H2O2 in the Absence of 0 2
so2 Henry's constant (M atm-l at 20 "C) gas-phase concn (ppb)
1.43' 30
NO2 O.Olb
50
HzOz 1.05 X 106 1
Aqueous S(1V) Oxidation Rates (M s-l) contribution at pH 4 3.4 x 10-8 3.5 x 10-6 3.4 x 10-7 3.5 x 10-6 contribution at pH 5 contribution at pH 6 3.6 X 10-6 3.5 X 10-8 a
Reference 34. Reference 35. Reference 32.
formation of a transient intermediate. Similarly, formation of a transient intermediate could also occur in the nitrogen dioxide-sulfite ion reaction, and the intermediate could rapidly form nitrite ion and sulfite radical. This would be consistent with the results that we observe. The reaction of NO2 with S(1V) was studied by Lee and Schwartz (6), using low concentrations of aqueous S(1V) and gaseous NOz. They did not observe dithionate as a product, although it may have been present below their detection limit. They observed a ratio of about 1.5 for nitrite ion produced to sulfite consumed in the absence of oxygen. This is slightly higher than the values obtained in our studies. It is consistent with a value of 2 for the branching ratio (4b/4a) of reaction 4. Ellison and Eckert (7) studied the effect of NO2 and Nz04 on the oxidation of S(1V) by 0 2 in the presence of Fe(I1) and Mn(I1) ions under acidic conditions in a batch system. They observed that NO2 enhanced the oxidation rate under some conditions. Rosenburg and Grotta (5) also observed that NO2 promoted the oxidation of S(1V). It should be pointed out that NO2 is not a catalyst for S(IV) oxidation, since it is not regenerated in the way that metal ions are. I t can produce sulfite radical, which can then initiate a radical oxidation chain involving reactions 8-10. The above mechanism for the reaction of nitrogen dioxide with sulfite ion suggests that this reaction could make a significant contribution to the oxidation of S(1V) in atmospheric aqueous aerosolsunder suitable conditions. Oxidation of aqueous S(1V)has been under investigation for many years (25-29). Hydrogen peroxide is generally considered to be the most important S(IV) oxidizing agent in the atmosphere (30). We can compare the relative rates of oxidation of S(1V) in aqueous aerosols by hydrogen peroxide and nitrogen dioxide using reaction rate expressions from the literature (9, 31-33). Using the Henry's constants and the gas-phase concentrations for SOZ,H202, and NOz listed in Table 111, the rates of oxidation of S(IV) in aqueous aerosols by NO2 and HzOz were calculated in the absence of 02. The gas-phase concentrations listed are what might be found in an urban atmosphere. Gasphase S(1V) oxidation is slow compared to oxidation in aqueous aerosols. The calculations were made for aqueous aerosols at pH 4,5, and 6, assuming solution concentrations in equilibrium with the gaseous concentrations listed. The calculated S(1V) oxidation rates are listed in Table 111. The contribution due to the direct reaction of NO2 with S(1V)is comparable to the contribution due to the H202S(1V)reaction only at the highest pH condition. However, the NOz-S(IV) reaction will generate SO3*-radicals that can continue to oxidize S(1V) via radical chain reactions. This process can effectively multiply the oxidation rate to many times the rate due to the N02-S(IV) reaction alone. 2166
Envlron. Scl. Technol., Vol. 27, No. IO, 1993
The experiments done in this study with low NOz concentrations in the presence of oxygen show that many S032- molecules can be oxidized per NO2 absorbed. Consequently, oxidation of S(1V) by NO2 could be comparable to oxidation by HZOZover a range of conditions. The oxidation of S(IV) by metal ions has not been included in this discussion because the process is quite complicated (28). The metal ion oxidation process may involve the SO3'- and SO5'- radicals and could interact with the NO2 oxidation process. The oxidation of S(1V) by NO2 could also play an important role in aqueous flue gas desulfurization systems, particularly in systems where NO is converted to NO2, such as the PhoSNOX process (36). The absorbed SO2 is often precipitated with calcium salts. It is desirable to oxidize the dissolved S(1V) before it precipitates, since calcium sulfite has a number of undesirable properties. The presence of NO2 would significantly enhance the S(1V) oxidation rate. A comprehensive model would be required to obtain a quantitative measure of the relative importance of the S(1V) oxidation pathways in atmospheric aerosols and is beyond the scope of this investigation. However, the contribution to atmospheric S(IV)oxidation from the NOZS(1V) reaction should not be neglected, particularly in polluted environments with relatively high concentrations of nitrogen dioxide.
Author Supplied Registry Numbers: NO2, 1010244-0; S032-,14265-45-3. Literature Cited (1) Nash, T. J. Chem. SOC.A 1970, 3023. (2) Takeuchi, H.; Ando, M.; Kizawa, N. Ind. Eng. Chem. Proc. Des. Dev. 1977, 16, 303. (3) Takeuchi, H.; Takahashi, K.; Kizawa, N. Ind. Eng. Chem. Proc. Des. Dev. 1977, 16, 486. (4) Sato, T. S.; Matani, S.; Okabe, T. Abstracts of Papers, Presented at the 177th National Meeting of the American Chemical Society, Honolulu, HI, April 1, 1979; American Chemical Society: Washington, DC, 1979; INDE-210. (5) Rosenberg, H. S.; Grotta, H. M. Environ. Sci. Technol. 1980, 14, 470. (6) Lee, Y. N.; Schwartz, S. E. In Precipitation Scauenging, Dry Deposition, and Resuspension; Pruppacher, H. R., Semonin, R. G., Slinn, W. G. N., Eds.; Elsevier: New York, 1983; pp 453-470. (7) Ellison, T. K.; Eckert C. A. J. Phys. Chem. 1984,88, 2335. ( 8 ) Huie, R. E.; Neta, P. J. Phys. Chem. 1986, 90, 1193. (9) Clifton, C. L.; Albtein, N.; Huie,R. E.Enuiron. Sci. Technol. 1988, 22, 586. (10) Lee, Y. N.; Schwartz, S. E. J. Phys. Chem. 1981, 85, 840. (11) Hall, T. C.; Blacet, F. E. J. Chem. Phys. 1952, 20, 1745. (12) Chang, S. G.; Littlejohn, D.; Lin, N. H. In Flue Gas Desulfurization; Hudson, J. L., Rochelle, G. T., Eds.; ACS Symposium Series 188; American Chemical Society: Washington, DC, 1982; pp 127-152. (13) Littlejohn, D.; Chang, S. G. Enuiron. Sci. Technol. 1984, 18, 305. 1972,94, (14) Hayon, E.; Teinin, A.; Wilf, J. J.Am. Chem. SOC. 47. (15) Eriksen, T. E. J . Chem. SOC.,Faraday Trans. 1 1974, 70, 208. (16) Eriksen, T. E. Radiochem. Radioanal. Lett. 1974,16, 147. (17) Waygood, S. J.; McElroy, W. J. J. Chem. SOC.,Faraday Trans. 1992,88, 1525. (18) Huie, R. E.; Clifton, C. L.; Altstein, N. Radiat. Phys. Chem. 1989, 33, 361. (19) Deister, U.; Warneck, P. J.Phys. Chem. 1990, 94, 2191. (20) Huie, R. E.; Neta, P. Atmos. Environ. 1987, 21, 1743.
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Received for review January 19, 1993. Revised manuscript received June 1, 1993. Accepted June 15, 1993.' Abstract published in Advance ACS Abstracts,August 15,1993.
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