I
H. F. JOHNSTONE and D. R. COUGHANOWR] University of Illinois, Urbana, 111.
Absorption of Sulfur Dioxide from Air Oxidation in Drops Containing Dissolved Catalysts SuLFuRic ACID aerosol, often, present in the air of industrial communities, is partly responsible for reduced visibility in smog (74, 20). Sulfur dioxide and its oxidation products have been held responsible in part for the disasters in the Meuse Valley (1930) ( 5 ) , Donora, Pa. (1948) (78), and London (1948 and 1952) (27). Pattle and others (2, 76, 77) concluded that sulfuric acid mist may be involved in the toxicity of contaminated fogs. Gerhard and Johnstone (7) studied the rate of photochemical oxidation of sulfur dioxide in moist air and found that droplets of sulfuric acid of 0.2- to 0.4micron diameter are formed in intense sunlight. Because of the slow rate of the reaction they suggest that oxidation of the gas in fog droplets may be important. Ellis (4)found that the sulfuric acid-sulfur dioxide ratio in a London fog is higher than in fog-free air. Iron and manganese salts are catalysts for the oxidation of sulfur dioxide in acid solutions (70). Dust particles containing these elements ( 7 3 , 74) may serve as nuclei for the formation of fog droplets. T h e rate of sulfur dioxide oxidation in drops of catalyst solutions suspended in air was studied and compared with the rate of sulfur dioxide absorption and oxidation by drops containing hydrogen peroxide.
dioxide of known concentration. T h e chamber was placed in a thermostat a t 25 f 0.2" C. and arranged so that the suspended drop could be moved in or out, observed under a 60x microscope equipped with a calibrated graticule, and its size measured. After exposure to the gas, the drop was removed from the fiber and diluted with water, and acid content determined in a 0.1-ml. Lucite microconductivity cell. As little as 1y of sulfuric acid could be determined accurately ( 3 ) . The partial pressure of sulfur dioxide in the gas was determined by passing a known volume of nitrogen through a sample of the solution with which the gas surrounding the drop was in equilibrium and absorbing the sulfur dioxide in dilute hydrogen peroxide. T h e concentration of the acid formed in the peroxide was determined by conductivity. This method was accurate and gave results consistent with other methods of vapor pressure determination. T h e equilibrium concentration of sulfur dioxide in the exposure chamber corresponded to 20 to 200 p.p.m. a t atmospheric pressure. T h e catalyst concentration was varied from 250 to 1000 p.p.m. and the drops measured with the microscope were from 700 to 900 microns in diameter.
Absorption b y Drops of Peroxide
For comparison with the catalyst solution, the rate of absorption of sulfur dioxide from air by stationary drops of dilute hydrogen peroxide was measured (Table I). For each drop the rate of absorption was constant and slightly higher than that predicted by the theory of mass transfer for diffusion to a sphere in a quiescent gas, for which NINushould be 2 (79). T h e reaction of sulfur dioxide and hydrogen peroxide is rapid, so that the equilibrium vapor pressure of the gas a t the surface of the drop should be zero; it is also highly exothermic. T h e abnormally high rates may have been caused by density currents set up by evaporation of water at the surface of the drops. Possibly, the areas of the elongated drops were larger than those of the spheres of the diameters observed under the microscope. Rate of Reaction in Solutions of Manganese Sulfate
To interpret the data on absorption and oxidation in the catalyst solution, the kinetics of the reaction of dissolved sulfur dioxide and oxygen was studied in solutions of manganese sulfate ( 3 ) , by rapidly mixing solutions of sulfur dioxide
Experimental
An attempt was made to use a suspension of droplets of dilute solutions in air to simulate a natural fog, but it was difficult to control relative humidity within narrow limits near saturation and collect the droplets after a fixed time of exposure. Finally, a microtechnique was developed by which a single drop of catalyst solution could be exposed to humid air containing sulfur dioxide and analyzed for its acid content. The smallest drop that could be handled conveniently was about 300 microns in diameter. This is larger than the droplets in natural fogs, but Garner and Skelland (6) show that there is no more circulation in the larger drops. The drops were suspended on a 15micron glass fiber in a n exposure chamber (Figure 1) in which the gas was in equilibrium with a solution of sulfur Present address, School of Chemical and Metallurgical Engineering, Purdue University, Lafayette, Ind.
Figure 1. An exposure chamber was used for determining sulfuric acid in drops in air
MICROSCOPE
CLASS DISCHARGE
r
- I
1
VOL. 50, NO. 8
AUOUST 1958
1169
I
C
C*
rFigure 2.
Concentration profiles of gas being absorbed by a drop and reacting Rate of Absorption and Oxidation by Drops Containing Catalyst
with solutions of oxygen containing manganese sulfate and quenching the reaction after fixed intervals of time by pipetting samples into standard iodine to determine the amount of sulfur dioxide remaining. I n some cases the catalyst solution was acidified by adding reagent sulfuric acid or acid made by bubbling sulfur dioxide and oxygen through the solution to simulate conditions in fog droplets nucleated by dust particles. The rate of reaction always remained constant until the oxygen content of the solution was essentially depleted. From the initial slope of the rate curves the apparent zero-order reaction constant, ko, expressed as millimoles/ (liter) (minute), was found. Table I1 shows that when sulfuric acid is not present in the solution initially the rate constant is proportional to the square of the concentration of manganese sulfate. The results agree with those of Hoather and Goodeve, who studied the catalyzed reaction a t 25" and 35" C. (8). When sulfuric acid was added to the catalyst solution before mixing, the initial rate decreased as the acid concentration increased. T h e reaction was easily inhibited by impurities. I n some cases there appeared to be an induction period, which Hoather and Goodeve attributed to the presence of oxidizable inhibitors. Inhibition by low concentrations of phenols has been noted (70), and by the organic accelerator vaporized by passing a n air-sulfur dioxide mixture through a short length of rubber tubing ( 7 ) .
Table 1.
From the analyses of the drops exposed to the sulfur dioxide-air atmosphere at saturated humidity for measured intervals of time the rate of absorption of the gas expressed as moles/ (min.)(sq. cm.) was found from the slopes of the straight lines of N D us.
DoB/D. R=-
dND
d(BD,/D)
(vol. of drop in liters) (surface area of drop)
I t is assumed that ko is approximately independent of acid concentration. Solving Equation 1, and substituting the boundary condition, c = ci at r = ro,
T h e shape of the concentration profile is shown in Figure 2. There is a critical concentration, c*, at the surface of the drop, for which the concentration at the center just reaches zero. (4)
When ci < c*, c becomes zero before the dissolved sulfur dioxide reaches the center of the drop, as shown at rl in Figure 2. Solution of Equation 2 is
Do is the diameter of a reference drop used so that the data could be analyzed in groups of drops of uniform size and composition; D is the average diameter of the drop at the beginning and end of the exposure. T h e drop size increased slightly when the concentration of acid was high, because of absorption of water vapor from the gas in equilibrium with the dilute solution of sulfur dioxide. Discussion of Results
I t is assumed that a steady state exists and that the liquid in the drop is quiescent, so that mass transfer takes place by molecular diffusion. The resistance to absorption is entirely within the drop and the concentration of the dissolved
When the dissolved sulfur dioxide penetrates only to r l , the reaction takes place in the spherical shell outside of rl; the rate of reaction, Q, which is also the rate of absorption of sulfur dioxide, is
When ci > c*, the sulfur dioxide penetrates to the center and the reaction takes place throughout the drop; the rate is at a maximum and equal to
From Equations 5, 6, and 7,
Absorption of Sulfur Dioxide b y Water Drops Containing Hydrogen Peroxide (Temperature, 25' C.) HzSO4
Diameter of Drop, Microns 738 876 638 867 695 780 778 855 697
Exposure Time,
SO2 Concn.
Concn.
Min.
in Gas, P.P.M.
Barometer,, Mm. Hg
5.0 4.9 6.6 8.8 6.9 10.9 22.0 12.2 12.6
197 197 89 89 89 89 20 20 19
748 748 746 746 743 743 733 733 745
R,
Moles/(Min.) in Drops, (Sq.Cm.) Mole/L. x 106 0.70 0.48 0.56 0.42 0.50 0.62 0.31 0.15 0.22
1.72 1.42 0.90 0.69 0.84 0.74 0.18 0.18 0.20
INDUSTRIAL AND ENGINEERING CHEMISTRY
Nusselt NN,o., Nu 2.32 2.29 2.36 2.42 2.39 2.36 2.63 2.74 2.66
Av.
1 170
gas at the surface is its solubility at the partial pressure in the surrounding gas, which, in the present case, depends on the acid concentration in the drop. For this work, sulfur dioxide was the limiting reactant, so that the problem may be treated as one of absorption and reaction of a single gas. The concentration of sulfur dioxide along any radius of the drop is expressed by the differential equation,
2.46
(9)
A simplified form of Equation 8 may be derived by using the binomial expansion of (1 - Q / Q 0 ) 2 / 3and neglecting the terms with exponents greater than 2. For Q/Qo + Q, or T I / T O + 1 the result is
For convenience, Equations 8 and 10 may be written in terms of the rate of absorption per unit area of drop surface,
SULFUR DIOXIDE IN A I R
R. When the definition of R and the value of c * are substituted in Equation 8,
2.0
I
I
I
I
l
l
1
1
I
1
I
1
I
I
l
l
'
?2 X
SLOPE-0.5
-
Equation 10 may be simplified to give
For small distances of penetration, where Q / Q Oapproaches zero,
T h e solubility of sulfur dioxide in dilute acid solutions was found from the equations and data given by Johnstone and Leppla ( 7 7 ) . For the conditions in the present work, the dissolved gas exists entirely as un-ionized molecules and the solubility follows Henry's law cu
=
HPso,
5 5 E-
0.1 lo
0.
Absorption by Drops of Other Catalysts Some measurements were m-ade on the rate of the absorption of sulfur dioxide
p.p.m.
Figure 3. Rates of absorption of sulfur dioxide from air by drops of manganese sulfate solution agree well with theory Measured rates for drops 7 5 2 microns in diameter
--- Fits data calculated from best value of ko, 2.6 mmoles/(liter)(min.)
(14)
At 25', the value of H is 1.23 moles/ (kg.) (atm.). Partial pressures of sulfur dioxide in equilibrium with several solutions of the gas in dilute sulfuric acid confirmed the calculated values. Experimental rates of absorption and oxidation b y drops of manganese sulfate solution are compared with those predicted by the theory in Figure 3. The slope of one half on the logarithmic scale for the parallel lines for 500 and 1000 p.p.m. agrees with Equation 12 for the case of small penetration into the drop before the solute is entirely consumed. T h e surface of the drop is assumed to be saturated with the sulfur dioxide in the gas. The value of D 1used is 2 x 10-5 sq. cm. per second, which is approximately the value given for sulfuric acid in water (79,p.\ 22). By extrapolation of the lines to 1 p.p.m. of sulfur dioxide in the gas, the values of R are 4.4 X and 7 . 2 X 20-6 millimole/(min.) (sq. cm.) and the corresponding values of ko would be 6.4 and 17.2 mmoles/(liter) (min.) for the two concentrations of the catalyst, respectively. At a lower concentration, 250 p.p.m. of manganese sulfate, the solute penetrates farther into the drop and Equation 11 applies. T h e agreement of the shapes of the curves with those predicted by the theory, and the consistency of the data, give confidence in the extrapolation of the results to smaller drops and lower concentrations of sulfur dioxide. Table I11 shows typical rates of absorption and distances of penetration of sulfur dioxide into the drops.
IO00
IO0
C O N C E N T R A T I O N OF SO2 I N CAS,
from air by drops containing ferrous and ferric sulfate (Table 111). The rate of oxidation of sulfur dioxide by drops containing iron sulfate was so much slower than that with manganese sulfate that as much as 50,000 p.p.m. of the catalyst had to be used to approach the rate for 250 p.p.m. of manganese sulfate. T h e use of such high concentrations caused some difficulties in determining the acid content of the drops. This was best done by calibrating the conduc-
Table
Run
II.
tivity cell with solutions of sulfuric acid containing ferric sulfate and measuring the conductivity of the diluted drop before hydrolysis and precipitation of ferric hydroxide. Addition of small amounts of copper sulfate to the manganese sulfate solution scarcely affected the rate of absorption by the drops, although 'similar ratios inhibited absorption and oxidation in a bubbler. This shows the effect of the induction period when a n inhibitor is present.
Reaction of Dissolved Sulfur Dioxide and Oxygen in Solutions of Manganese Sulfate Initial 'CO Initial Mmoles/ HzSOa MnSOa, Concn., Oaa Concn., Temp., (L)(Min.) P.P.M. Moles/L. Mmoles/L. O c. x 100
12-2 12-3 12-6 12-7 12-8 12-4 6-5 6-6 11-2 12-1 12-9 13-2 13-3 14-1 14-2 15-1 15-2 15-3 15-4 15-5 16-1 16-2 16-3
4.0 4.0 5.0 5.0 5.0 8.0 10.0. 10.0 10.0 10.0 12.5 12.5 12.5 12.5 12.5 12.5 12.5 12.5 12.5 12.5 12.5 12.5 12.5
12-5
14.0
0 0 0
0.1od
2.14 2.10 2.52 2.01 2.03 1.96 1.94 1.89 2.03 2.16 2.10 1.95 2.06 2.04 2.06 1.96 1.91 2.04 2.09 2.01 1.95 1.93 2.02
1.07 0.78 1.69 0.84 0.44 0.69 0.53 0.58 0.72 0.86 0.66 1.09 0.62 1.30 1.16 0.83 0.99 1.18 1.30 1.65 1.19 1.20 1.76
0
2.22
1.06
0 0
0 0 0 0
0 0 0.002c 0.002= 0.02 0.02e 0:035d 0.002d 0.002d 0.02d 0* 05d 0.002d 0.02d
... ... ...
...
24.9 24.9 24.9 24.9 24.9 24.9 24.7 24.7 24.9 24.9 24.9 24.9 24.9 24.6 25.1 24.9 24.9 24.9 24.8 24.8 24.8 24.8 24.9
2.6 2.6 2.7 3.1 3.8 11.1 15.3 15.7 17.0 17.1 20.1 5.9 6.9 2.5 2.6 1.8 8.2 7.8 3.3 0.7 6.7 3.3 0.5
0.58
24.8
32.6
0.66
0.80 0.63 0.71 0.66 0.66 0.65 0.72
... ... ..I
I
.
.
0.56 0.46 0.43 0.40
...
0.38 0.37
Calculated from amount of sulfur dioxide oxidized. Concentration when reaction stopped. Acid produced by reaction of sulfur dioxide and oxygen in solutions of manganese sulfate. Reagent acid.
VOL. 50, NO. 8
AUGUST 1958
1171
There was no measurable rate of absorption by drops containing 1000 p.p.m. of nickel sulfate when sulfur dioxide in the gas was 350 p.p,m.
droplets formed by the photochemical oxidation, the residual acid droplets formed by the catalytic reaction in fog are probably larger than those formed by photochemical oxidation.
Conclusions
A cknowledgrnent T h e authors acknowledge with pleasure financial support of this work by the American Petroleum Institute as Project SF-9 of the Smoke and Fumes Committee. Nomenclature = concentration of dissolved solute
in drop, moles per cc. = concentration of solute a t surface
of drop, moles per cc. concentration of un-ionized H2SOa, moles per kg. HzO = minimum concentration of solute a t surface of drop for which solute penetrates to center, moles per cc. = diameter of drop, cm. or microns = diameter of standard reference drop = diffusivity of solute in liquid, sq. cm. per minute = Henry’s law constant, moles/ (kg. HzO)(atm.) = reaction rate constant for zero order reaction, millimoles/ (min.) (1.) = distance of penetration of solute into the drop, cm. or microns = concentration of H2S04in drop, moles per liter = Nusselt group for mass transfer, dimensionless = partial pressure of SO2. atm. = rate of absorption by drop, moles per minute = maximum rate of absorption by drop, moles per min. = radial distance from center of drop. cm. = radius of drop, cm. = radial distance from center of drop at which concentration of solute becomes zero, cm. =
Rate of Absorption and Oxidation of Sulfur Dioxide by Drops of Catalyst Solutions Rate of SO2
Catalyst MnSOa FedSOds FeSO1 MnSOa Fe2(S04)3
+
MnS04
+ CuSOa
Catalyst, Concn., P.P.M. 250 500 1,000 53,200 5,000 500 2,660 500 5 1,000
Absorp-
Concn. tion, Diam. a t Drop Moles/ Coiicn. of Surface (Min.) Penetraof S O Z , Drop, l\.Ioles/L. (Sq. Cm.) tion,O P.P.M. Microns X IO‘ X lo7 Microns 298 120 197 295 360 120
700
250
780
700 700 760 790 770
NiSO4 Calculated from Eq. 13, using Dz of 2 X small distances of penetration, approximate value.
1 172
8
rate of absorption per unit area of drop surface. moles/’(min.) (sq. cm.) = exposure time, minutes
=
literature Cited
T h e rate of formation of sulfuric acid in a drop of catalyst solution suspended in air containing sulfur dioxide depends on the partial pressure of sulfur dioxide and the nature and concentration of the catalyst. When the drop contains dilute manganese sulfate the rate of acid formation is on the order of one tenth of the rate in drops of hydrogen peroxide and agrees with a theory for reaction throughout the drop. For high concentrations of the catalyst, the reaction takes place in the outer shell of the drop and the sulfur dioxide does not penetrate to the center. While the composition and number of dust particles in air vary greatly, the possible rate of acid formation in a fog nucleated by manganese salts may be estimated from the data by assuming that each droplet is formed on a 1-micron crystal of manganese sulfate, or manganese oxide that is converted to the sulfate, that the diameter of the droplets is 20 microns, and that the water content of the fog is 0.2 gram per cubic meter, or 4.8 X l o 7 droplets per cubic meter. These assumptions are consistent with published information on fogs and clouds (9, 72, 75). I n a n atmosphere of 1 p.p.m. of sulfur dioxide the estimated rate of oxidation is 1% per minute, nearly 500 times the rate of the photochemical reaction found by Gerhard and Johnstone for intense sunlight. T h e catalytic reaction takes place in the dark as well as in light. Hoivever, as the concentration of sulfuric acid in a n acid droplet increases with the inverse cube of the diameter as the droplet evaporates in air of low relative humidity, and the initial fog droplets are one or two orders of magnitude larger than the
Table 111.
R
350
700 sq.
INDUSTRIAL AND ENGINEERING CHEMISTRY
3.7 1.5 2.5
...
0.32 0.53 1 .o 0.13 0.07 0.45
...
0.46
...
0
... ...
cm./sec.
227 72 60
...
...
...
... ...
Bassett, H., Parker, W. G., J. Chenz. Soc. 1951, p. 1540. Collumbine, H., Pattle, K. E., Burgess, F.. “Toxicity of Fog,” 7th International Congress of Comparative Pathology. 1955. Coughanowr, D. R., Ph.D. thesis, Vniversity of Illinois, 1956 (Vniversitv Microfilms. Inc.. .Ann Arbor, h&h.). Ellis: B. A,, “Investigation of Atmospheric Pollution,’‘ 17th Report, Dept. Ind. Sci. Research (Brit.), D. 38.1931. Fir’ket, ‘ J., Trans. Faradaj Soc. 32, 1192 (1936). Garner, F. H., Skelland, A. H. P., Trans. Inst. Chem. Engrs. (London)
29, 315 (1951). Gerhard, E. R., Johnstone, H. F., IND.ENG.CHEM.47, 972 (1955). Hoather, R. C., Goodel-e, C. F., Trans. Faiaday Soc. 30, 1149 (I 19343. Houehton, H . G., Radford, W. EL, PGers Phys. Oceanog. Meteorol., M a s s . Inst. Technol. and Woods Hole Oceanogl-afiiz. Inst. 6 , KO. 4 (1938).
(10) Johnstone, H. F.,I h D . ENG.CI-IEM. 23, 559 (1931). (11) Johnstone, I-I. F.: Leppla, P. W., J . Am. Chem. Soc. 56, 2233 (1934). (12) Kampe, H. J.. in “American Institute of Physics Handbook,’’ 1957 ed., pp. 2-134, lfcGraw-Hill, New York. (13) Katz, M., Am. Ind. Hyg. Assoc. Quart. 13, 4 (1952). (14) Los Angeles County Air Pollution Control District, “Teclinical and Administrative Report on .Air Pollution in Los Angeles County,’’ p. 16, 1949-50. (15) Neiburger, M., Wurtele, A4. G.. Chem. Revs. 44, 321 (1949). (16) Pattle, R. E., Burgess, F., Collumbine, H., .J. Pathol. Bactei-id. 22, 219 (1956). (17) Pattle, R. E., Collumbine, H., Brit. M e d . J . 2, 913 (1956). (18) Schrenck, H. H., Heimann, H., Clayton, G. D., Gefafer, W. M . , L‘. S. Pub. Health Bull. 306, 162 (1949). (19) Sherwood, T. K., Pigford, R. L., “Absorption and Extraction,” 2nd ed., p. 72, McGraw-Hill, New York, 1952. (20) Stanford Research Institute, “Smog Problem in Los Angeles County,” pp. 101-9, 1954. (21) JVilkens, E. T., Smokeless Aw 24, No. 88, 89 (1 953). RECEIVED for review September 7, 1957 A4CCEPTF,D .April 11, 1958
LIT, 0.79b 0.21 0.17
... ...
... ...
...
As Eq. 13 applies only for
Divisions of Industrial and Engineering and Water, Sewage, and Sanitation Chemistry, Symposium on Air Pollution, 132nd Meeting, ACS, Sew York, N. Y., September 1957. Material supplementary to this article has been deposited a s Document KO. 5571 with the AD1 Auxiliary Publications Project, Photoduplication Service, Library of Congress, Washington 25, D. C . A copy may be secured by citing the document number and by remitting $1.25 for photoprints or $1.25 for 35-mm. microfilm. Advance payment is required. Make checks or money orders payable to Chief, Photoduplication Service; Library of Congress.
