the regression line of the device on the continuous monitor was calculated. The results of these statistical tests for all results, and for all results less those obtained in tests 6 and 7 , are shown in Table IV. The significance tests on the regression analysis indicate that the value of the gradient obtained is very significant, and that of the intercept not significant a t all. As a further check on the significance of the results, a t test was carried out on all the results, using the difference between each pair of readings and the null hypothesis that the differences do not differ significantly from zero. A value of 1.15was found for t . Entering the t table a t u = 16, P is found to lie between 0.25 and 0.30 and thus our null hypothesis is upheld.
Table IV. Statistical Comparison of Results Parameter Correlation coefficient Gradient 95% confidence limits t statistic Probability U
Intercept 95% confidence limits t statistic Probability U
All results
Omitting 6 and 7
0.995 1.02 f0.06 39 -65 % 16 i-30 f141 0.453
0.995 1.02 f0.06 35 -75% 14
((0.1 %
((0.1 %
16
14
4-25 f169 0.313
Conclusions
Acknowledgment
Reiszner and West claim that their permeation device will provide reasonable mean concentrations of sulfur dioxide without the use of power for sampling, a t low cost and with the minimum of analysis. The results of the tests described above generally support this claim. The device offers not only an alternative to the lead dioxide candle as a power-free sulfur dioxide monitor but can be expected in most circumstances to provide results that are better than merely indicative. Furthermore, depending on the level of sulfur dioxide in the air surrounding the device, it has been shown that results can be obtained over periods as short as 1 h or as long as 19 days; the upper time limit is dependent only on the loss of water from the device by evaporation. The chemistry is practically specific for sulfur dioxide. The toxicity of the reagent could be a drawback in use, particularly at unattended measuring sites.
The author is indebted to Dr. Egan, the Government Chemist, London, for the opportunity to test one of Professor West’s permeation devices. L i t e r a t u r e Cited (1) Reiszner, K. D., West,
P. W., Enuiron. Sei. Technol., 7, 526-32 (1973). (2) West, P. W., Gaeke, G. C., Anal. Chem., 28,1816-19 (1956). (3) Littlewood, A,, J . Sei. Instrum., 44,878-80 (1967). (4) Killick, C. M., ibid., ( J . Phys. E ) , Series 2,2, 1017-20 (1969). (5) British Standards Institution, London, “Methods for the Measurement of Air Pollution: The Lead Dioxide Method”, BS1747: Part 4, 1969. (6) Saltzman, B. E., Anal. Chem.. 33. 1100-12 (1961). (7) Liang, S.F., Sternling, C. V., Galloway, T. R., J.A.P.C.A., 23, 605-7 (1973). Received for review February 25, 1975. Accepted January 5 , 1976.
Reactivity of Zinc Oxide Fume with Sulfur Dioxide in Air William L. Dyson and James E. Ouon* Environmental Health Engineering, Northwestern University, Evanston, 111. 6020 1
The reactions of metal oxide particles with sulfur dioxide have many implications for understanding the air pollution phenomenon and its control. The formation of acids and sulfates under certain conditions are well known. In urban atmospheres, zinc is as prevalent as lead and the average concentration of zinc is 0.7 llg/m3 ( I ) . Zinc ammonium sulfate was considered to be partly responsible for the irritant properties of the fog in the Donora episode ( 2 ) . Amdur and Corn ( 3 ) reported that zinc sulfate and zinc ammonium sulfate aerosols were respiratory irritants capa’ ble of producing significant increases in pulmonary flow resistance in guinea pigs. The synergistic health effects of some combinations of gaseous and aerosol pollutants are generally recognized. In studying the reaction between sulfur dioxide and zinc oxide as a possible air cleaning process, Gressingh et al. ( 4 ) found that up to 50% by weight of sulfur dioxide could be absorbed by zinc oxide. They implied that zinc sulfite was the main reaction product. The health implications of zinc sulfite are unknown. Hence, depending on the nature of the reaction product and the reactivity, the sulfur dioxide-zinc oxide reaction may be a secondary source of an irritant in the atmosphere or an atmospheric sink for sulfur dioxide. The focus of this 476
Environmental Science & Technology
study was to evaluate the reactivity or capacity of zinc oxide fume for reaction with sulfur dioxide in air. The concentrations of sulfur dioxide used were 4.0-17.6 ppm. The reactivity was evaluated a t temperatures of 15, 2 5 , and 35 “C and for relative humidities of 2-95%. Differentiation was made between sulfite and sulfate reaction products. Some information on the order of the reaction and the initial reaction rate was also obtained. S h r i n k i n g Core M o d e l
If, as suggested by Gressingh et al. ( 4 ) ,zinc sulfite is the main product for the reaction of zinc oxide with sulfur dioxide in air, and the reaction does not proceed in the absence of water vapor, then the reaction falls within the general class of noncatalytic gas solid reactions. Relationships for the conversion of zinc oxide as a function of exposure time when any one interaction step is rate controlling have been derived by Wen ( 5 ) for a constant diffusing gaseous reactant concentration in the gas mixture. For the case where gas film diffusion controls, the relationship is
The reactivity of zinc oxide fume with sulfur dioxide in air was studied to determine its dependence on sulfur dioxide concentration, temperature, and humidity. Zinc oxide fume was exposed to about 10 ppm of sulfur dioxide a t 15, 25, and 35 "C and 2-95% relative humidity. Hydrated zinc sulfite was the main reaction product. The reactivity was found to be 110 pg S032-/mg ZnO a t 25 "C, 50% relative humidity, and sulfur dioxide concentrations between 4.0
and 17.6 ppm. At 25 "C and 91% relative humidity, the reactivity increased to 760 pg S0S2-/mg ZnO. The reactivity increased with decreasing temperature a t all humidity levels. Sharp increases in the reactivity occurred a t humidity levels corresponding to the equilibrium phase transition of zinc oxide to zinc hydroxide. The reaction was limited by diffusion of water vapor through the sulfite product layer.
For the case where diffusion through the sulfite product layer controls, it is
the geometric mean diameter, was 3.87 X a t 25 "C, indicating laminar flow. The system temperature was maintained a t 25 & 1 "C by room air for the majority of runs. For exposures a t 35 "C, the elevated temperature was obtained by immersing the mixers in a constant temperature bath a t 50 "C, using flexible heating tape on the line from the last mixer to the filter reactor, and heating the filter reactor with infrared lamps. High humidities a t 35 f 2 "C could be attained, but control was cumbersome. For exposures a t 15 "C, the entire system, including the water bubblers, all mixers, and the filter reactor, was immersed in a constant temperature bath. Immediately after exposure, each zinc oxide filter was weighed to determine the weight increase during exposure. Dependent upon the weight gain, the filter was then immersed in 50 or 100 ml of 0.04M potassium tetrachloromercurate (TCM) solution in a 600-ml beaker. Analyses for sulfite and sulfate were made using portions of this solution. Both sulfur dioxide and sulfite were determined using Method A of the modified West and Gaeke procedure as described by Scaringelli et al. (8). Sulfur dioxide in the gas mixture was absorbed into 10 ml of TCM solution in a midget impinger. A flow of 1.0 l./min was sampled for 0.52.0 min. The entire content of the impinger was used for analysis. For sulfite analysis from 1-10 ml of the TCM solution was used, depending on the expected amount of sulfite as estimated by the filter weight gain. Colorimetric determination for sulfite was made a t 548 nm (Spectronic 20, Bausch and Lomb, Inc.). Sulfate was determined by a modified nephelometric method (9) using a Hach Turbidimeter (Model 1860) and 20-ml aliquots of TCM solution. The sensitivities of the analytical methods are 0.5 pg SO', 3.0-6.0 pug s03'-, and 5 pg SO4*- for the quantities of TCM solution used. For sulfur dioxide and sulfite measurements, the accuracy was about 0.75 pg SO2 or SO:,'- for sample concentration of 25 pg SO2 or SOz2- per 10 ml of TCM solution. The accuracy was in the order of 2-3% for all concentrations of sulfur dioxide or sulfite measured. For sulfate measurements the accuracy was only 3.5 pg SO4'for a sample concentration of 10 pg sod2- per 20 ml of TCM solution. The accuracy improved with higher sulfate concentrations and worsened with lower sulfate concentrations. Measurements of sulfite and sulfate were made for several types of blanks. The results (Supplementary Material, Table I) indicated that no correction was necessary for the sulfite or sulfate determinations.
and for small X ,
For the case where are chemical reaction controls and a single diffusing gaseous reactant of constant concentration, the relationship is
These equations may be used to determine the rate-controlling step for a reaction from plots o f t vs. X and log-log plots o f t vs. .' Experimental Apparatus and Procedures Zinc oxide particles were generated by the decomposition of zinc acetate aerosols (6). Dry air a t 5.6 l./min was passed through an atomizer containing a 10% zinc acetate solution. The atomizer effluent containing zinc acetate aerosols was passed into a Pyrex tube (2.5 cm i.d., 60 cm long) heated by an electric furnace to a centerline temperature of 340 & 10 "C. Cooling air a t 6.4 l./min was mixed with the particle suspension a t the exit of the furnace and the entire flow was passed through a 5-1. flask from which samples were taken. Tared Teflon filters (Mitex LSWP, Millipore Corp.) were supported in the filter reactor by a Teflon-coated screen. Sampling for about 3.5 min a t 9.0 l./min deposited about 7.0 mg of zinc oxide particles in a thin layer over a 45.6-cm2 area of the filter. Because it was suspected that water and various organic decomposition products were adsorbed on the surface of the particles, the filters with deposited particles were dried overnight a t 110 "C. This reduced the amount of collected material an average of 30% to about 5.0 mg. X-ray diffraction confirmed that the particles were zinc oxide. The "dried" filters with deposited zinc oxide were placed in Petri dishes and stored in a desiccator until needed. A particle size analysis was made from enlargements of electron microscope photomicrographs to 20 OOOX. The geometric mean diameter and geometric standard deviation of the zinc oxide particles were 0.14 pm and 2.72, respectively. Gas mixtures for exposure of the zinc oxide filters were generated in essentially the same way described by Chun and Quon ( 7 ) . The only difference was that sulfur dioxide was premixed with a dry nitrogen stream (50 ml/min) because the asbestos plug flow meter had a tendency to condense water when exposed to high humidity conditions for long periods, thus affecting the sulfur dioxide flow. The interstitial gas velocity through the filter during exposure was 0.44 cm/s and the particle Reynolds number, based on
Results and Discussion Measurements of the sulfate portion of the reaction products indicated that the sulfate was a very minor component when compared with the sulfite content (Supplementary Material, Table 11). The average sulfate content for 26 measurements a t various exposure times and conditions was 8 pg S0d2-/mg ZnO. Thus, rate and reactivity measurements are based on sulfite content. Volume 10, Number 5, May 1976
477
Table 1. Initial Reaction Rate Constant at High Sulfur Dioxide Concentrations (Temperature, 25OC; N ~ , , 3 . 8 7X Experimental series and sample
ZnO sample, mg
SOP concn,
L1 2 3 4 5 6 7 8 9 10
4.6 4.9 4.4 4.6 5.3 5.5 4.5 3.9 4.3 4.8
28 150 145 149 98 53 152 105 148 150
8.0
I
Relative Water vapor humidity, YO pressure, mol/cm3
PPm
87 87 50 50 50 50 90 90 70 70
11.3 X lo-' 6.5 X lo-'
11.7 X lo-' 9.1 x 10-8
I
I
ZnS03,2&H,0
-
s'ZnSO3.2 H,O
7.0
,' ,' ,
,'
6.0-
# ,
z 9
0
I
I
P
I ,#'
, ,
I
,
I
,
,,.g I
I
5.0I
,
,
,,
Zn(HSO&
,,
,
Exposure time, min
Weight gain, mg
Sulfite formed, fig S032-(mg ZnO)-'. min-'
Rate constant, cm*min-' (moil cm3)-'
2 3 3 3 3 3 3 3 3 3
0.70 1.20 0.43 0.33 0.39 0.34 1.44 0.96 0.66 0.67
61.0 65.2 26.0 19.1 19.6 16.5 85.2 65.6 40.9 37.2
77 82 101 74 76 64 100 77 81 73
In each case the sulfite content was limited initially by the sulfur dioxide input a t the concentration used. Upon continued exposure the sulfite content increased until the reactivity was reached. This reactivity was approximately 110 pg S032-/mg ZnO for the different sulfur dioxide concentrations used. The reactivity measurement was fairly reproducible as shown by a standard deviation of 20 pg S032-/ mg ZnO for 12 measurements. The range of the measurements, 80-137 pg S032-/mg ZnO, may be due to slight differences in the zinc oxide fume sample and variations in the water vapor pressures. Another set of filters (Series F) was exposed to 10.9 ppm of sulfur dioxide at 25 "C and 91% relative humidity and compared with exposures a t 12.2 ppm and 50% relative hu-
SERIES E
0
I20
I:O
Figure 1. Correlation of
2:O 30 410 5:O SULFITE CONTENT, mq
60
7b
SO, = 17.6 ppm
zinc oxide filter weight gain with sulfite con-
T = 25OC R H = 50%
tent
Three main reaction products containing sulfite ions are possible for the sulfur dioxide-zinc oxide reaction in the presence of water vapor. They are zinc sulfite hemipentahydrate, ZnS03.21/2HzO, the dihydrate, ZnS03-2H20, and zinc bisulfite, Zn(HS03)z. Figure 1 shows a comparison of the filter weight gain during exposure with the total quantity of sulfite measured for that filter. Various exposure conditions are represented, but sulfite contents greater than 1.0 mg are from exposures at high humidities. The weight gains correspond closely to the expected weight gain for zinc sulfite dihydrate as the reaction product. Although it is likely that the sulfite reaction products are a combination of the hydrates and the bisulfite in proportions that apparently indicate the dihydrate as the product, calculations of initial reaction rate were simply based on zinc sulfite dihydrate as the reaction product. When exposed filters were immersed in distilled water rather than TCM solution, the resulting pH was 6.4-7.3. This indicated that the reaction products contained little or no free acid. Sulfite formation as a function of time for exposures to 4.0, 12.2, and 17.6 ppm sulfur dioxide at 25 "C and 50% relative humidity (Series C, D, and E) are shown in Figure 2. 478
Environmental Science & Technology
SERIES D
140II
CI
$
1
120-
0
-1
I-z
1
801
W
L
z 401 w
t LL
20
3
0
m
-
RH = '
'
1
'
'
'
'
'
'
'
'
50% '
'
'
-
SERIES C
140120
-
A
100-
80 60 40
A
-
A A
SO, * 4.0 ppm T = 25-C R H = 50%
i
20- A 0
0
"
40
"
80
"
"
"
120 160 200 EXPOSURE TIME, min
"
240
"
280
Figure 2. Filter sulfite content variation with sulfur dioxide concentration and exposure time at 50% relative humidity
0 0 O
€a 0
~ 0
" 60
'
"
SO, = 12.2 pprn RH = 50%
' ~ ' 240 300 EXPOSURE TIME, min
120
~
SERIES D
180
~ 360
"
420
"
Figure 3. Filter sulfite content as a function of exposure time at 12 ppm sulfur dioxide and two different relative himidities
midity (Series D) (Figure 3). The reactivity a t 91% relative humidity was approximately 760 pg S032-/mg ZnO. This represents a conversion of 77% of the available zinc oxide t o zinc sulfite. Variation of the reactivity with water vapor pressure was further examined by exposing zinc oxide fume to sulfur dioxide a t relative humidities from 8-91% a t 25 "C (Series G). The exposure time necessary for reactivity measurements was different a t various humidity levels. The necessary exposure time was determined by measuring the filter weight gain periodically during exposure and continuing the exposure until no further weight increase was observed. The results of these exposures (Figure 4) show the tremendous effect of water vapor on the sulfur dioxide-zinc oxide reaction. The reactivity ranged from 14-806 pg S0S2-/mg ZnO. When the water vapor concentration was low, little or no zinc sulfite was formed. This collaborates the conclusion of Gressingh et al. ( 4 ) that water vapor was essential for reaction. When the water vapor pressure approached saturation a t 25 "C (23.8 mm Hg), practically all of the zinc oxide was converted to zinc sulfite with continued exposure to sulfur dioxide. The data a t 25 "C in Figure 4 also show an order of magnitude increase in the slope of the reactivity curve a t higher humidities with the transition occurring a t about 1 7 mm Hg water vapor pressure. Although several possible explanations of the effect of the water vapor were tested, the one which best correlates with the observations a t high humidities involves the equilibrium between zinc oxides and water and the formation of zinc hydroxide. A phase diagram of the zinc oxide-water system is shown in Figure 5 ( 1 0 ) .The calculated equilibrium relationships are considered to be accurate only for a temperature variation of *IO "C. They do, however, show the general trend that can be expected. Following the 25 "C isotherm for increasing water vapor concentrations, one finds a phase transformation to zinc hydroxide a t approximately 18 mm Hg. This corresponds closely to the water vapor pressure a t which the transition in the slope of the reactivity occurred a t 25 "C (Figure 4). Two series (H and I) of exposures a t 12 ppm sulfur diox-
~
ide were made a t 35 O C and 15 "C respectively to see if a change in the slope of the reactivity curves could be observed a t these temperatures. From the zinc oxide-water phase diagram, a transition would be expected a t about 8 and 38 mm Hg water vapor pressure a t 15 and 35 "C respectively. The results of these exposures are also shown in Figure 4. The transition can be seen in the data obtained a t 15 "C (Series I). Although the transition is less distinct than for the data a t 25 "C, i t does correspond to the predicted occurrence a t about 8 mm Hg water vapor pressure. Again, for vapor pressures near saturation a t the temperature used (12.8 mm Hg a t 15 "C), practically all of the zinc oxide was converted to zinc sulfite. The data obtained a t 35 "C (Series H ) do not show a transition, although water vapor pressures up to 38 mm Hg were used. This may possibly be explained by uncertainty in the equilibrium diagram (Figure 5). Limitations of the experimental apparatus restricted the use of higher water vapor concentrations. When higher concentrations were attempted, condensation occurred within the system and on the filter. The sulfite determinations obtained could not be interpreted because of the possible absorption of sulfur dioxide by the condensed water. If water vapor concentrations approaching saturation could have been maintained, an upward turn in the reactivity curve would have been expected. Sulfite measurements obtained a t 15, 25, and 35 "C (Series I, G, and H) indicate that for a given humidity level, more sulfite is formed a t lower temperatures. Lower temperatures enhanced the reactivity a t all humidity levels, but especially when saturation for a given temperature was approached and the transition of zinc oxide to zinc hydroxide occurred. This corresponds to the general trend for physical adsorption where the quantity of vapor adsorbed decreases as the temperature increases. I t also suggests that the effect of water vapor is to convert the zinc oxide to
0 c
N
900
-
800
-
700SATURATION
0
E
\
'b"
600-
v)
m
-70
+ 0
-'C 15 25 35
rnrnHg
12.8
24.8 42.2
a
-60 W
B
X
500-
-50
IW I-
z
z
$
N
400'
-40
0
c z
w
c
5 300-
0 W
-30
ul 2
200
-
100
-
a
- 20
VAPOR PRESSURE OF WATER, rnm Hg
Figure 4. Variation of maximum filter sulfite content (reactivity) with temperature and water vapor pressure Volume 10, Number 5, May 1976
479
Table II. Effective Diffusivities When Diffusion of Water Vapor Through Product Layer Controls Reaction Rate Exptl serles
I"
s
D(H20)', cm2/min
flnltl.l,
1.38X IO-" 3.32 X lo-" 3.65X IO-" E 2.65 X F 8.18x 10-10 K a D calculated from Equation 3.
C D
60
E E K W I-
Relative humidity, %
50-
C..I, mln
mln
60 15 15 120 60
50 50 50 91 98
90 30 30 240 120
LL
0
40-
3
v
W
)
-
a K
g 3
30-
2i
8.
"'I 0.30
IO
0 1 , l 110
a
a n a
g
'
20 I ' 30 I ' TEMPERATURE,
I40
"
"50
60
OC
Flgure 5. Equilibrium phase diagram for zinc oxide-water system
g
0.10 SERIES F
0.08
SOz = 10.9ppm
a
W K
T = 25.C
0 IL 0.05b
the zinc hydroxide rather than convert the sulfur dioxide to sulfurous acid prior to the reaction. Equations 1 to 4 for the shrinking core model assume the reaction to be limited by a critical diffusing gaseous component. In our case, the diffusion of water vapor was the limiting component. This is shown by plots of p vs. t (e.g., Figure 6) for Series C, D, E, F, and K and by the initial sulfite formation rate for different combinations of sulfur dioxide and water vapor (Series L, Table I). A plot of p vs. t for Series F is shown in Figure 6. The extent of conversion of zinc oxide based on complete reaction of the input sulfur dioxide and the amount of zinc oxide present are indicated by the dashed line. The actual extent of conversion of zinc oxide deviates from the sulfur dioxide input-limiting case a t exposure times between 60 and 120 min. During the next period, between 120 and 240 min exposure time, the slope of the conversion curve is approximately one-half, indicating that diffusion of the critical gaseous reactant through the sulfite product layer was the rate controlling step (Equation 3). After 240 min exposure, very little conversion occurs and the slope approaches zero. The chemical reaction step was also found not to be the controlling step for the other experimental series (C, D, E, and K). For the period in each experimental series during which diffusion of water vapor through the sulfite product layer was controlling, the effective diffusivity was calculated from Equation 3 (Table 11). The effective diffusivity varied from 1.4 X lo-" to 82 X 10-l' cm2/min. These small values of diffusivity are indicative of a tight crust structure. For such cases, small structural changes in the crust can result in large changes in diffusivity. At high relative humidities, the large quantity of reaction product formed and possibly additional hydration of the crust material may lead to some rupturing of the crust, resulting in a large increase in diffusivity. Ten zinc oxide filters were exposed to high sulfur dioxide concentrations in an attempt to overcome the limitation of 480
Environmental Science & Technology
W
1
0.02
I
20
30 40
60 80 100
,
I
200 300 EXPOSURE TIME, min
, 1 , 1
500
I
I
800
Figure 6. Variation of the extent of reaction parameter with exposure time (Series F)
low sulfur dioxide input rates. The objective was to obtain initial reaction rate data which was not limited by the input rate. The maximum possible sulfur dioxide output of the experimental apparatus was used. The sulfite produced after three minutes of exposure was calculated from the filter weight gain, assuming the product to be ZnS03.2H20. Weight gains of approximately 10 pg could be estimated (Mettler H 16 balance). Only exposures a t relative humidities of 50% or more were made since the weight gain a t low humidities would have been too small to measure with any accuracy. The initial rate constants for these 10 exposures a t high sulfur dioxide concentrations are shown in Table 11. Variation of the sulfur dioxide concentration from 50-150 ppm a t 50% relative humidity changed the sulfite formation rate very little (Filters 3 to 6). However, increasing the relative humidity from 50-90% a t 150 ppm sulfur dioxide increased the sulfite formation rate from about 23 to about 75 pg SO$min-larng Zn0-l (Filters 3, 4, and 7-10). These results show that the initial reaction rate is independent of the sulfur dioxide concentration and approximately second order with respect to the concentration of water vapor. The reaction rate constant calculated on this basis varied from 64101 with an average of 80 crn-min-l (mol/cm3)-'. This agreement of the reaction rate constant is good considering
the sensitivity of the weighings. The rate constant calculated from the data of the Gressingh et al. ( 4 ) for 15 min exposure was 1 2 cmarnin-l (mol/cm3)-l. For several of the above filters the exposure was continued beyond the initial 3 min, and additional weighings were made. Analysis by the methods described previously indicated that the reaction becomes a t least partially controlled by diffusion through the product layer very quickly. Loglog plots of cp vs. t had slopes between one and one-half beginning with the data point a t 3 min. This suggests that the reaction is controlled predominantly by diffusion of the critical gaseous reactant through the sulfite product layer after the initial few minutes. This is one possible explanation of why the initial rate constant calculated from the 15-min data point of Gressingh et al. ( 4 ) was much lower than the initial rate constant from the present exposures. Another reason may be that the experimental arrangement used by Gressingh et al. ( 4 ) was a porous bed and does not fit the theoretical model used here. Gas transfer and reaction opportunities are dependent upon the particle size, interstitial velocity of gas flow past the particle, and particle Reynolds number. The values of these parameters used in this study are well within the expected range of values encountered in the atmosphere (7). However, one cannot extrapolate results obtained in our study directly to the atmosphere since many other variables may be involved. Nevertheless, the reactivity of zinc oxide shows that for an average urban atmosphere containing 0.7 pg Zn/m3 ( I ) up to 0.8 pg S0s2-/m3 may potentially be formed, depending upon the chemical form of the zinc compound, the composition of the atmospheric aerosol, and other atmospheric parameters. These experimental results have other implications to the field of air pollution. When low temperatures and high humidities are present, zinc oxide has definite potential as an absorber of sulfur dioxide ( 4 ) . Also, these results indicate reactions and reaction products which may be found in zinc smelter plumes since sulfur dioxide and zinc oxide are emitted simultaneously. Finally, where zinc is used as a galvanizing material or as zinc oxide in white paint, the reaction of zinc oxide with sulfur dioxide may be of economic importance. Nomenclature
C = concentration of fluid in the gas mixture, mol/l. or mol/cm3 D = effective diffusivity of critical gaseous reactant through zinc sulfite product layer, cm2/min k = reaction rate constant k , = mass transfer coefficient for diffusion of gaseous reactant through gas film, cm/min M = molecular weight of zinc oxide, g/mol
R
= geometric mean radius of zinc oxide particle prior to reaction, cm t = exposure time, min X = fraction of zinc oxide converted to zinc sulfite, dimensionless Greek Letters p = density of zinc oxide, g/cm3 cp = extent of reaction parameter = Subscripts
1 - (1 - X)1/3
f = diffusing gaseous reactant x = order of reaction for critical diffusing reactant Superscript
x = order of reaction for critical diffusing reactant
L i t e r a t u r e Cited (1) US. Department of Health, Edgcation and Welfare, Public Health Service, Cincinnati, Ohio, “Air Quality Data from the National Air Sampling Networks and Contributing State and Local Networks”, 1966. (2) Hemeon, W. C. L., A M A Arch. Ind. Health, 11, 397-402 ( 1955). (3) Amdur, M. O., Corn, M., A m . Ind. Hyg. Assoc. J., 24, 326-33 (1963). (4) Gressingh, L. E., Graefe, A. F., Miller, F. E., Barber, H., “Applicability of Aqueous Solutions to the Removal of SO2 from Flue Gases”, Final Rep., Vol. I, Aeroject General Corp, El Monte, Calif., Contract P H 86-68-77,1970;NTIS: P B 196780. (5) Wen. C. Y.. Znd. E m . Chem.. 60 (9). 34-54 (1968). (6) Marshall, B. S., Telford, I:, Wood, R., Analyst, 96 (1145), 569-78 (1971). ( 7 ) Chun, K. C., Quon, J. E., Enuiron. Sci. Technol., 7 (6), 532-8 (1973). (8) Scaringelli, F. P., Saltzman, B. E., Frey, S. A., Anal. Chem., 39, 1709-19 (1967). (9) Keily, H. J., Rodgers, L. B., ibid., 27,759-64 (1955). (10) “Gmelin’s Handbuch der Anorganischen Chemie”, Vol. 32, Zink, p 935, 1956.
Received for review August 9, 1974. Resubmitted September 15, 1975. Accepted December 10, 1975. This study was supported, i n part, by funds from the Department of Civil Engineering, Northwestern University, Euanston, Ill., 60201; and the Bureau of Community and Environmental Management, ECA, U S P H S , Research Training Grant No. 5-TOI-EC-00014. Mention of commercial products is for identification only and does not constitute endorsement by a n y branch of the U. s. Government. Supplementary Material Available. One page of Supplementary Table I (Chemical Analyses of Filter Blanks) and 4 pages of Supplementary Table I1 (Measurements for Reaction of Zinc Oxide and Sulfur Dioxide in Air) will appear following these pages in the microfilm edition of this volume of the journal. Photocopies of the supplementary material from this paper only or microfiche (105 X 148 mm, 24X reduction negatives) containing all of the supplementary material for the papers in this issue may be obtained from the Business Operations Office, Journal Department, American Chemical Society, 1155 16th St., N.W., Washington, D.C. 20036. Remit check or money order for $4.00 for photocopy or $2.50 for microfiche, referring to code number ES&T-76-476.
Volume 10, Number 5, May 1976
481
