CURRENT RESEARCH Effects of Relative Humidity and Temperature on IronCatalyzed Oxidation of SO, in Atmospheric Aerosols Johnny Freiberg' Department of Chemistry, T h e Johns Hopkins University, Baltimore, Md. 21 218
2 ) considered the iron-catalyzed oxidation of atmospheric SO2 to acid mist to be a major cause of production of s04'- in urban atmospheres. The roles of relative humidity and temperature, in promoting this process, are quantitatively evaluated in connection with the mechanism of SO2 oxidation to HzSO4 and its subsequent neutralization by ambient "3. Thus, the oxidation rate iincreases rapidly with increasing relative humidity (particularly a t high relative humidities), and it decreases by about an order of magnitude for an increment of 5°C in temperature. There is a qualitative agreement between these results and the observed rates of SO2 oxidation in the atmosphere. W Previous investigators (1,
Among the most troublesome of air pollutants are SO2 and its oxidation products. Sulfur dioxide and sulfuric acid aerosols have been held responsible for damage to vegetation and materials (3, 4 ) and for harmful effects on humans ( 5 ) . Although sulfur dioxide is monitored daily in most industrial urbanized areas, few investigations were reported on S O Z oxidation rates in urban atmosphere (6). Sulfur dioxide is removed from the atmosphere by photooxidation in the gas phase and catalytic oxidation in cloud droplets. thereby transforming the gas to particulates containing S 0 4 2 - . Photochemical oxidation predominates in dry areas and has a rate which gives a residence time for SOz of 21 to 42 days (7). The catalytic oxidation predominates in humid industrial areas (such as in England and the U.S.) and has a rate which gives an average residence times of 5 days, although in England it is 8.75 hr (7). Few attempts have been made to measure the oxidation rate of SO2 in solutions as dilute as those occurring in water droplets in the atmosphere. Until Junge and Ryan's work ( I ) , the chemical literature contained no quantitative information on the rates of oxidation of SO2 to HzSO4 in such dilute solutions. Since then, experimental data significant to the heterogeneous catalyzed oxidation of atmospheric SO2 were gathered (8-11). Likenise, the photochemical oxidation of either SO2 or of a mixture of different hydrocarbons, SOz, NO, and ozone was studied intensively, Only a few of the above investigations were concerned with the effects of the temperature, T , and the relative humidity (RH) on the SO2 oxidation in the atmosphere and their results seem to conflict. In irradiated mixtures of 2-methyl-2 butene, S O , and SO:!, increasing RH increases the S 0 4 2 aerosol formation (12). The same effect was reported by
Johnstone and Coughanowr (91, and Johnstone and Moll (IO) who investigated the catalytic oxidation of SO2 in the presence of various metal ions. The aerosol formation by irradiation of 2-pentene, NOz, and SO2 was not affected in any way by RH (13).Wilson and Levy (14) reported that increasing RH decreases aerosol formation in a 1-butene-NO-SO2 mixture. The same effect was reported by Harkins and Nicksic (15) who irradiated auto exhaust from a high-sulfur fuel and by Groblicki and Nebel ( 1 6 ) who irradiated a propylene-nitric oxide-SO2 system. With respect to the effect of T on the oxidation rate, Harkins and Nicksic (15) found a threefold decrease in aerosol formation when the T was increased from 50°F to 100"F, while Groblicki and Nebel data indicated that 40% less aerosol is formed when T is decreased from 95°F to
80°F. Neytzell de Wilde and Taverner (8) found that the ironcatalyzed oxidation of SO2 in solution is accelerated by temperature, and Hoather and Goodeve ( 1 7 ) , in their work on SO2 oxidation catalyzed by manganese, found the ratio of reaction velocities to be 2.09 for a change of 5°C in the range 30-35°C.
Rate of Iron-Catalyzed Oxidation of SO2 These disagreements about the effects of the RH and T on the reaction rate of SO2 oxidation may be due to the different mechanisms by which SO2 is oxidized to S42-. However, none of the aforementioned investigators presented a mechanism for their system of SO2 oxidation. Recently Freiberg (18) deduced the mechanism, rate expression. and the rate constant for the iron-catalyzed oxidation of SO2 in acid solutions; a t low [Fe3-1, characteristic for the droplet solutions in the atmosphere ( 1j. this rate expression may be written
d[SO,'-] -- K,K,'[H?SOJII,?[Fe'+],, (1) dt [H+I,, where KO is a rate constant, Ki is the first dissociation constant of sulfurous acid; [HzSOs], , [Fe3f],, and [Ht], are concentrations in the droplet solution. (Square brackets [ ] are used to describe concentrations of the indicated chemical species.) The rate of SO2 consumption (disappearance) in the gas phase depends on the rate of s04'- production in the liquid droplet (as in Equation 1) and on the volume of water droplets where the S 0 4 2 - production occurs. Thus, one can write
Present address, Aeronautical Research Associates of Princeton, Inc , Princeton, N . J . 08540 Volume 8 , Number 8 , August 1974
731
where [Wj is the volume of water droplets per unit volume of air and [Fe3-] [Fe3+lu[Wl. Sulfur dioxide obeys Henry's law in dilute acidic solutions when dissociation is accounted for (19);
KH[HLSOJI,,
= PSO?
as
F.\H
(3)
where Pho2is the partial pressure of SO2 in air and K H is the Henry constant. The Henry constant has been determined between 0-50°C by Johnstone and Leppla (19). The expression for SO2 solubility (Equation 3) will be coupled with the rate of reaction in the liquid phase (Equation 2) as follows: Henry's expression (Equation 3) can be rewritten, since [SO21 in plumes is low enough that it obeys the ideal gas law
and
1
(12)
Droplet growth, subsequent to the production of s04'-, can be described in terms of the relationship between solute concentrations in the droplets and the vapor pressure of the surrounding atmosphere. It is difficult to evaluate this relationship accurately, since the solution in the droplet contains two solutes, (NH4)2S04 and Fe3-, both of which are ionized and may interact. The vapor pressure lowering coefficient of Fe3+ used is smaller than that of (?;H4)2S04, and [Fe3-] in the droplets is much smaller than [(NH4)2S04]. Therefore the contribution of Fe3+ to the vapor pressure lowering in the droplet is very small and may be neglected by taking the vapor pressure lowering coefficient as equal to the effect of (NH4)2S04 alone:
(4) where ps is the Ostwald constant, [H2S03IiLis the concentration of sulfurous acid in the liquid phase, and [SO21 is the concentration of sulfur dioxide in the gas phase. Substitution of Equation 4 into Equation 2 yields
The rate expression (Equation 5) is strongly dependent on pH so that the SO2 oxidized to H2SO4 will decrease the oxidation rate and eventually stop the reaction. As demonstrated by Junge and Ryan ( I ) , there is sufficient ammonia in the ambient atmosphere, especially in industrial areas, to neutralize much of the sulfuric acid formed and maintain the pH a t buffered values where reaction can occur a t reasonable rates. Furthermore, measurements of [H+] in aerosol particles a t four locations in Europe indicated that sulfuric acid buffered by ammonia is the cause of acidity in aerosol particles (20). Neutralization of HzSO4 by absorption of NH3 from air proceeds almost instantaneously (21). Therefore, one can write an equilibrium relationship between H' and HzS04 neutralized on one side and the ammonia on the other. Combining the ionization and equilibrium expressions:
one gets:
where F\H BnKn[NH3]/Ku. Assuming 100%dissociation for HzS04, one writes:
[H']
+ [h",']
=
2[S04'-]
(10)
Substitution of Equation 9 into Equation 10 yields:
[H+][l
+ F \ H ]= 2[So,'-]
and by rearrangement: 732
Environmental S c i e n c e & Technology
(11)
where Xz is the pressure-lowering coefficient for (NH4)2S04 and p / p o , the ratio of vapor pressure t o that a t saturation a t the same T. is the relative humidity ( p / p o = RH). Moreover, in Equation 13 the effect which the curvature of the droplet has on the pressure lowering was neglected. This is a valid assumption for droplets which have a radius larger than 0.1 pm. The substitution of Equation 13 into Equation 12 yields:
Thus the pH will depend on [S042-], which varies with RH on the [NHsJ in the atmosphere, as well as on the dissociation constants of ammonia and water and the Ostwald constant for ammonia. These are all dependent on temperature, so that pH is dependent on RH as well as on
T. Now the rate of SO2 consumption (disappearance) may be expressed exclusively on parameters characteristic of the gas phase. The substitution of Equation 14 into Equation 5 yields:
- d[SO?I dt
where [SOzJ, [Fe3f], [NH3] are local concentrations in the gas phase. In the derivation of Equation 15, the values for [H2S03J and [H'] were determined through the equilibrium Equations 4 and 14, respectively, by the local environmental and water vapor. Also, in the concentration of S02, "3, development, the number and size of the water droplets were not taken into account. This can be true only if the rate of diffusion of SO2 to and in the droplet is fast with respect to the rate of oxidation. Diffusion in the droplet was considered to be the limiting process in the oxidation of SO2 catalyzed by MnS04 in artificial droplets of -700 pm diameter and a t saturated humidity (9). However, Johnstone and Coughanowr did not compare the rate of diffusion with the rate of oxidation because they did not have an expression for the latter. In this study, the oxidation of SO2 catalyzed by iron is considered in aerosol droplets at lower RH and in droplets which are expected to be smaller than 10 pm in diameter.
Thus, the oxidation rate should be smaller because the iron is a poorer catalyst than manganese (9) and because the relative humidities in the atmosphere are usually lower than the saturation. Additionally, the diffusion time is smaller (owing to the smaller droplets) than the corresponding diffusion time in Johnstone and Coughanowr’s droplets. Moreover, the reaction rate expression of ironcatalyzed oxidation of S O z is available (18). Now, one can evaluate both the rate of diffusion in a droplet and the rate of oxidation and ascertain how good the approximation of neglecting the size and the number of droplets is. The computations made by Freiberg ( 1 8 ) show that the characteristic time of reaction is greater than the characteristic time of diff’usion in the droplet for most atmospheric conditions. Another interesting point is that at the same temperature, an increase in reaction rate due to an increase in R H will increase the diffusion time while a commensurate increase in reaction rate due to an increase in [”SI will not affect the diffusion time as much. This is so because in the former case more water will condense and the size of droplets will increase. Thus, neglect of the number and size of the droplets should be a better approximation when the reaction rate increases due to increase in [NH3] than when it increases due to increase in RH. For this work, wt: shall retain these assumptions, and treat Equation 15 as the local rate of depletion of SO2 in any small average volume. Dependence ofReaction Rate on Relative HumiditJ, The relative humidity is the atmospheric parameter considered to have the greatest influence on the catalytic oxidation of SO2 in solutions. Gartrell et al. (22), Stephens and McCaldin (23),and Berger et al. (24) observed that high RH promotes the oxidation of SOz; however. none of these investigators tried to express mathematically the dependence of the extent of conversion on RH. In aerosol droplets, sulfur dioxide and 0 2 diffuse throughout the solution and react catalytically in the presence of iron salts forming more HzS04. The formation of HzSO4 and its subsequent neutralization to (NH4)2S04 lowers the vapor pressure and induces more water to condense. continuing the process. Equation 13 shows that more water condenses for the same amounts of solute in solution as the RH increases. For example, when the RH increases from 80Yc to 90% the amount of water condensed increases twofold, from 80% to 95% it increases fourfold, and from 80% to 99% it increases twentyfold. The condensation of more water in the droplets affects the reaction rate of‘ SO2 oxidation in three ways: It increases the amount of soluble SO2 available for oxidation; it increases the pH on which the oxidation rate is strongly dependent, and it dilutes [Fe-3]. The first and the third effects compensate one another, such that the RH affects the oxidation rate only by changing the pH. Because [H-] 1 - RH (from Equation 14) and the rate of SO2 oxidation d[SOz]/dt = 1/[H+I3(from Equation 15), the net result is d[SOz]/dt = 1/[1 - RHI3 (as written in Equation 15). Hence, the rate of iron-catalyzed oxidation of SOZ is “cubically” dependent on RH. For example, when the RH increases from 80% to 90%, the rate of oxidation increases eightfold; from 8070 t o 9 5 7 ~RH, it increases sixty-four-fold. J
Dependence of Oxidation Rate ofSO2 on Temperature Previous investigations have not discussed how the temperature influences the rate of catalytic oxidation of SO2 in the atmosphere. Ordinarily one expects the reaction rate to increase with T because the rate constant for the
Table I 533n30,2K0Kn3
KO,
Ks,
T, ~m3 mol “ C m o l m i n md
5 10 15 20 25 30
6.45 9.10 11.20 15.60 20.40 26.30
20.9 18.4 16.6 14.5 13.0 11.8
pa
60.0 51.0 42.9 35.5 30.0 26.3
P,
3560 2780 2180 1780 1480 1160
K,,, mol ii7
K, X 108,
0.0148 0.0157 0.0165 0.0171 0.0177 0.0182
1.85 2.91 4.41 6.91 10.06 14.70
mol
(m:)
2
Kw3 x 10-36 m’ x m x
‘
2349.743 269.813 29.130 3.501 0.559 0.075
catalytic oxidation of SO2 increases with T . There are, however, other physical-chemical processes which influence the reaction in a different way as a function of temperature. The total effect is a marked decrease in the yield of the reaction as T increases. This can be understood as follows: The way in which the physical-chemical processes depend upon temperature is governed by the dependence of the reaction rate constants, the dissociation constants, and solubility constants on temperature. The factor K,2P,3/3,2KoKn3/KLL3from Equation 15 contains the dependence on T By inspection of Table I, one can see that the rate constants for catalytic oxidation of SOz in solution ( K O )and the dissociation constant of amproduce an increase in yield with T. monia in water (Kn) on the other hand, the Ostwald constants of NH3(Pn), of SOZ(P,).and the dissociation constants HzSO3(K,) and of HzO(K,) cause a decrease in rate with T A close inspection of the temperature dependence of each factor shows that the second group of parameters dominates the determination of the overall temperature dependence. The net result is a decrease in reaction rate of iron-catalyzed oxidation of SO2 by about an order of magnitude for an increment of 5°C. The substantial decrease in rate of the SO2 oxidation in atmospheric aerosols with increasing T is a new result which has not been described previously in the literature. It is, however, in accord with the fact that serious air pollution episodes have occurred during late fall and early winter months, when T was low and RH was high. Thus, the air pollution incidents a t Donora. Pa. (October 1948); Meuse Valley, Belgium (January 1911. December 1930, and December 1935); London, England (December 1952); and in worldnide episodes (Kovember to December 1962) took place a t low T in stagnant weather conditions at high RH. Conclusions The rate of iron-catalyzed oxidation of SO2 is strongly dependent on RH. particularly at high RH. This is due primarily to the pH in droplets that increases with RH. Likewise, the oxidation rate of SO2 in droplets is influenced by a number of physical-chemical processes which depend on temperature differently. The overall result is a decrease in the reaction yield with a temperature increase. The results agree qualitatively with the observed rates of SO2 oxidation in the atmosphere. Acknou~ledgment The author is indebted to Glenn R. Hilst and Everett Thiele, who kindly gave good advice. and to Arlene Spears for her typing. V o l u m e 8 . N u m b e r 8 , August 1974
733
Literature Cited (1) Junge, C. E., Ryan. T . G., Quart. J . Roy. Meteorol. SOC.,84, 46 (1958). (2) Foster, P. M., Atmos. Environ.. 3, 157 (1969). (3) Thomas, M . D., “Air Pollution,” W.H.O. Monograph Series 46, p p 233-75, Columbia University Press, New York, N.Y., 1961. (4) Burdick, L. R., U.S. Bureau of Mines, Inform. Cir. 7064. 1939. (5) Amdur, M . O., h t . J. Air Poilut., 1,170 (1959). (6) Urone, P., Enriron. Sei. Technol., 3,436 (1967). (7) Junge, C. E., J. Geophys. Res., 65,229 (1960). (8) Neytzell de Wilde, F. G., Taverner, L.. Second UK Int. Conf. on Peaceful Uses of Atomic Energy, Proc., 3,303 (1958). (9) Johnstone, H. F., Coughanowr, D. R., Ind. Enp. Chem., 50, 1169 (1958). (10) Johnstone, H . F., Moll, A. .J., ibid., 52,861 (1960). (11) Karraker, D . G., J . Phys. Chem., 67,871 (1963). (12) Schuck, E. A.. Doyle. G. J., Air Pollut. Found. Rep., p 29, Los Angeles, 1959. (13) Prager. M . .J., Stephens, E . R., Scott, W . E., Ind. Enp. Chem., 6,521 (1960). (14) Wilson, W. E.. Levy. A. J., J. Air Pollut. Contr. Ass., 20, 385 (1970).
(15) Harkins, J., Nicksic, S . W., ibid., 15,218 (1965). (16) Groblicki, P. J., Nebel, G . J., “Chemical Reactions in Urban Atmospheres,’’ C. S. Tuesday, E d . , p 263, American Elsevier Co., 1971. (17) Hoather, R. C., Goodeve. C. F., Trans. Faraday SOC.,30, 626,630,1149 (1934). (18) Freiberg, J., “The Catalytic Oxidation of SO? to Acid Mist in Dispersing Plumes,” P h D thesis. The Johns Hopkins University, 1972. (19) Johnstone, H. F . , Leppla, P. W., J. Amer. Chem. Soc. 56, 2233 (1934). (20) Junge, C., Schleich, G., Atmos. Enuiron., 5 , 165 (1971). (21) Cadle, R. D.. Robbins, R. C . , J . Phys. Chem.. 62,469 (1938). (22) Gartrell, F . E.. Thomas F. W., Carpenter, S. B., Amer. Ind. Hyg. Assoc. J . , 24, 113 (1963). (23) Stephens, E. T . , McCaldin, R . , Enciron. Sei. Technol.. 5, 615 (1971). (24) Berger, A. W , , Billings, C. E., e t al.. “Study of Reactions of Sulfur in Stack Plumes,” GCA Corp., First Annual Report GCA-TR-68-19G. 1969.
Received for revieu, October 12, 1973. Accepted March 28, 1974
Arsenic in Water by Flameless Atomic Absorption Spectrophotometry Kai C. Tam Freshwater Institute, Department of the Environment, 501 University Crescent, Winnipeg, Man., Canada R3T 2N6
.Arsenic in water is extracted with diethylammonium diethyldithiocarbamate in carbon tetrachloride and determined by atomic absorption spectrophotometry using the carbon rod atomizer. The method will determine arsenate, arsenite, and any organoarsenic compounds soluble in carbon tetrachloride. By using ultraviolet photooxidation to decompose organoarsenic compounds, the method determines total arsenic. No matrix interference is observed. Precision is f 0 . 4 gg/l. at 3.1 gg/l. and the detection limit is 1gg/l.
the rather short wavelength (1937 A) of the most sensitive arsenic resonance line. Although this can partially be overcome by using other types of flame, there seems to be more advantage with flameless techniques, such as thermal decomposition of arsine (12) or graphite furnace (23). Some preliminary work here confirmed the latter authors’ observation that matrix effects could be troublesome, requiring calibration by the standard additions method, for analysis of natural water. However, if arsenic is separated by a chelate-extraction method, interferences are avoided and a useful increase of sensitivity can be obtained.
Instruments Chronic poisoning has been reported to be caused by the utilization of drinking water containing 0.21-10.0 mg/l. of arsenic (As) (1-3). Standards for maximum allowable arsenic concentrations have been established by many agencies. The World Health Organization in 1958 ( 4 ) set a permissible limit of 0.2 mg/l. In 1963 the limit was revised to 0.05 mg/l. The United States Public Health Service ( 5 ) has a recommended acceptable concentration of 0.01 mg/l. and a maximum permissible limit of 0.05 mg/l. The Canadian Joint Committee on Drinking Water Standard (6) recommends the same levels. Inorganic arsenic in water is usually analyzed by evolution of arsine (AsH3) and colorimetric determination with silver diethyldithiocarbamate (7, 8) or by molybdenum blue methods, such as that of Johnson ( 9 ) , which require correction for any phosphate present. Kopp and Kroner (10) used emission spectrography. Neutron activation was used by Smales and Pate (11) for seawater. Other techniques such as X-ray fluorescence, polarography, and atomic absorption spectrophotometry (AAS) are potentially useful. There is difficulty with AAS methods if the air-acetylene flame is used because of flame absorption a t 734
Environmental Science & Technology
Means of irradiating samples in fused silica tubes (110ml capacity) a t 3-4 cm from a 450- or 550-W mercury arc lamp A Varian Techtron atomic absorption spectrophotometer (Model AA-5) equipped with carbon rod atomizer (Model 61 or Model 63) and strip chart recorder (Model A-25) using the following settings: AA-5
Wavelength, 1937 A Lamp current, 7 m A
Slit width, 300 p Slit height, 3 mm
C R A 6 1 or 63 Nitrogen flow, 1 I./min (CRA 61) or 4 I./rnin (CRA 63) Water flow. 0.5 I.lmin
Voltage setting Time, sec Range, 5 mV
Dry
Ash
3.5 20
6.25 15
Recorder
Atomize (step mode)
6.75
2
Chart speed, 25 in./hr
