Air pollution studies
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Photochemical Oxidation of Sulfur Dioxide in Air EARL R. GERHARD' AND H.
F. JOHNSTONE
University of Illinois, Urbana, 111.
The rate of the photochemical oxidation of sulfur dioxide in air was determined in the concentration range 5 to 30 p.p.m. The reaction was carried out in a Lucite cylinder in which the reaction mixture was exposed to ultraviolet light in the range from 3650 to 2950 A. The sulfuric acid drops formed as a result of the reaction were collected by a microtechnique involving impaction on an indicator film. The reaction is first order with respect to sulfur dioxide, and the rate is unaffected by the humidity, the presence or absence of salt nuclei, or by the concentration of nitrogen dioxide, in the ranges studied. The rate in artificial sunlight was in the order of 0.68% per hour and the mass median diameter of the particles was 0.19 to 0.45 microns. A few measurements in natural sunlight indicated the reaction rate to be approximately 0.1 % per hour. The gas phase oxidation of sulfur dioxide of naturally polluted atmospheres in natural sunlight would not be significant in effectively reducing the visibility, since it would require approximately 100 hours of intense noon sunlight to reduce the visibility to 1 mile. Other types of oxidation, such as the liquid phase reaction in fog droplets, or a combination of gas phase and liquid phase reactions at high humidity, may be more effective in reducing the visibility. These are being investigated.
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ULFUR dioxide in air undergoes a slow reaction in sunlight to form sulfuric acid and thus contributes t o haze and reduced visibility. The purpose of this work was to determine the rate of photochemical oxidation a t very low concentrations. The reaction conditions were similar to those in the atmosphere in industrial cities with sufficient water vapor present to permit aerosol or fog formation. The rate of aerosol formation and the effect of traces of certain impurities were also ascertained. The order of the reaction was determined by varying the sulfur dioxide concentration and the reaction time. The first observation of the role of sunlight in atmospheric pollution was made by Aitken a t the turn of the century ( 1 ) . He observed that fogs were produced over industrial cities by sunlight, and that the same result could be produced in the laboratory if a trace of sulfur dioxide was added t o air in a glass or quartz flask and exposed to sunlight. In order for this to occur, sulfur dioxide must absorb ultraviolet or visible light in the sun's spectrum. D a t a on solar radiation reaching the surface of the earth are shown in Figure 1. These are from the Smithsonian astrophysical measurements made a t Washington, D. C., as tabulated by Forsythe (6). Although the data around 3000 A. are of somewhat lower accuracy than that in the visible region, all investigators agree that there is a sharp cutoff between 2900 and 3000 A. that is attributed to the presence of ozone in the upper atmosphere ( 5 ) . The strongest sulfur dioxide absorption line occurs a t 1950 A., but the gas also absorbs in the region from 2400 to 3200 A. and slightly to 4000 A. (13). Values of the absorption coefficient k for pure gaseous sulfur dioxide a t 0" C. based on measurements a t 760 mm. mercury (10) are also shown in Figure 1. Although some decrease in the coefficient is to be expected a t higher tem1 Present address, Department of Chemical Engineering, University of Louisville, Louisville, Ky.
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peratures, sulfur dioxide does absorb solar radiation in the region from 2900 to 3300 A. The decomposition of pure sulfur dioxide in the near ultraviolet region was studied by Hill (9) who observed the formation of elemental sulfur. He concluded that 94% of the decomposition was due to the 3130 A. line, More recent photochemical studies of sulfur dioxide-oxygen mixtures have been reported by Blacet and Hall (8)who obtained a quantum yield for the decomposition of l x 10-8 from the 3130 A. line. With sulfur dioxide-oxygen mixtures only sulfuric acid was obtained. Since radiation in the range 2900 t o 3300 A. does not have sufficient energy to decompose the sulfur dioxide molecule it is probably activated. A possible reaction mechanism has been proposed by Blacet ( 2 ) . Experiments in Artificial Sunlight
Although it was desirable to obtain quantitative reaction data in sunlight, the effect of many variables could be studied more easily in a steady source of ultraviolet light approximating the spectrum and intensity of sunlight. General Electric sun lamp, Type RS, 275 watt, with a reflector bulb was used as the light source, The principal mercury lines obtained from this lamp are shown in Figure 1. The important lines corresponding to solar radiation which is absorbed by sulfur dioxide are a t 2967, 3022, and 3131 A. A diagram of the experimental equipment is shown in Figure 2. The air was passed through a bed of activated carbon to remove acidic and basic gases, which were present in the room air in sufficient quantities to affect the experimental results. Type CC-6 high efficiency cellulose-asbestos paper filter obtained from Arthur D. Little, Inc., was used to remove atmospheric aerosols. The relative humidity of the air was controlled by injection of steam.
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AIR POLLUTION Provision was made for the addition of salt nuclei by passing a part of the air stream over hot sodium chloride. Sulfur dioxide-air mixtures of 10 to 30 p.p.m. were made by supplying the gas from a cylinder a t rates of 20 t o 100 cc. per minute, through a calibrated glass capillary flowmeter, 10 inches long and 0.5 mm. inside diameter. For concentrations between 1 and 5 p.p.m., the gas was displaced by mercury from a 1-liter flask. The mercury was in turn displaced by oil from a constant speed proportioning gear pump.
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tions presented a problem. At 1 p.p.m., 1 liter of air contains only 2.85 micrograms of sulfur dioxide. Even if 100% conversion of the sulfur dioxide were obtained, 1 liter would contain only 4.37 micrograms of the acid. I n order to determine such small quantities, a special analytical technique based on the color change of an acid-base indicator film was developed. This technique, which is described in another paper (7), permitted the determination of less than 0.1 microgram of sulfuric acid. Experimental Procedure. At the beginning of each test, filteied air was passed through the reaction chamber for 15 t o 20 minutes. Continuous samples of the air were drawn through the automatic sulfur dioxide analyzer t o determine the blank. Steam was injected in the inlet air and regulated t o give the desired relative humidity in the reaction chamber. The sulfur dioxide concentrations at the inlet and outlet t o t h e reaction chamber were recorded continuously until they were equal and constant within 0.1 p.p.m If sodium chloride nuclei were to be added, 1% of the air was passed through the nuclei generator with the voltage across the heating element set to give a temperature of approximately 600 C. In some experiments, 1 to 5 p.p.m. of nitrogen dioxide were added from a I-liter flask by displacement with mercury. The reaction chamber was then closed and irradiation by the sun lamp was started. The reaction time was varied from 15 minutes to 4 hours. The temperature in the reaction chamber was recorded a t the end of the irradiation. The inlet line from the filtered air stream was opened a t the bottom of the chamber. From 200 to 1000 cc. samples of the aerosol were removed from the center of the chamber in from 10 to 60 seconds and impacted a t sonic velocity on the thymol blue indicator film. Particle size distribution of the aerosol was found in some runs by sampling a t several velocities. For more dilute aerosols, 25 liters of filtered air were passed through the chamber while the entire contents of the chamber were removed. The quantity of acid collected on the indicator film was determined by a colorimetric procedure using a photocell and microscope. I n this way runs were made with sulfur dioxide in air, with and without the addition of salt nuclei, at varying humidity, and with the addition of nitrogen dioxide to the sulfur dioxide-air mixture. Results. The d a t a obtained for sulfur dioxide-air mixtures containing 5, 10, 20, and 30 p.p.m. are shown in Figure 4. In these runs the reactor temperature varied between 75' and 89" F. and the relative humidity was varied from 32 t o 91 %. Duplicate runs showed a maximum deviation of 20% from the average.
The gas mixtures were analyzed by a Thomas autometer ( 1 6 ) . The reaction chamber was constructed of Lucite tubing, 13.8 cm. in diameter and 53.4 em. high, having a volume of 7.94 liters. The top of the chamber was 1/16-inchPlexiglas cemented t o the Lucite. The bottom of the chamber was a flanged Lucite plate sealed with a rubber gasket coated with a thin film of Lucite. This permitted removal of the sulfuric acid drops which deposited on the walls of the vessel by rinsing with water. The gas inlet and outlet tubes were made of saran. Plexiglas, Lucite, and saran were used because tests showed t h a t these materials did not adsorb the sulfur dioxide from dilute air-sulfur dioxide mixtures whereas glass and rubber tubing adsorbed appreciable quantities. Glass absorbed most of the ultraviolet in sunlight, while Lucite and Plexiglas ( inch thick) transmitted most of the ultraviolet above 2950 A. The transmission curve for Plexiglas Type-IA is shown in Figure 1. A sampling tube was placed in the center of the reaction chamber. The aerosol sample was drawn through a DUST single high velocity impactor by a vacuum pump capable of delivering 1 cubic foot of free air per minute at a vacuum of 20 inches of mercury. The rate of flow was determined by a calibrated jet and was controlled by a needle valve in combination with a AI R bleed line. The general details of the construction of the jet impactor and round jets and their calibrat,ion are described by Rane and Wong (Id). By use of a 0.353-mm. diameter round jet and sonic velocity in the throat it was possible to remove approximately 100% of the sulfuric acid particles 0.2 micron TO A S P I R A T O R in diameter and larger, and 50% of the 0.11-micron particles. The collection efficiency of the round jets is shown in Figure 3. Analysis of the small quantities of PUMP sulfuric acid obtained when working with such low sulfur dioxide concentraFigure 2. Schematic diagram of equipment
May 1955
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ENGINEERING, DESIGN, A N D PROCESS DEVELOPMENT The sulfur dioxide concentrations used were within 10% of the nominal concentration and no correction was made for deviations from t h e exact values.
Reaction I s First Order with Respect to Sulfur Dioxide Within the limits of error in the data, there is an approximately linear increase in sulfuric acid aerosol with time for a given sulfur dioxide Concentration. The per cent conversion per hour is essentially constant and averages 0.68, based on the data a t 10 and 20 p.p.m. During all the runs the sulfur dioxide concentration was essentially constant since less than 2% was converted to sulfuric acid. The rate of the reaction in the dark was negligible under these conditions. Even at 115 p.p.m. and 12 hours reaction time, a conversion of only 4 X per hour was obtained.
obtained in the absence of nitrogen dioxide. Apparently the ozone, which is produced catalytically by nitrogen dioxide ( S ) , is not significant in producing sulfur trioxide under these conditions. The results show that the addition of nitrogen dioxide in concentrations of 5 t o 20y0 of the sulfur dioxide concentration had no measurable effect on the reaction rate.
Effect of Salt Nuclei and Relative Humidity of Air Twelve runs were carried out in which sodium chloride nuclei were added t o the air stream. I n general, there was no effect of nucleation on the amount of aerosol produced. A series of runs was made at 10 p.p.m. sulfur dioxide for a reaction time of 60 minutes in which relative humidities of 32, 46, 71, and 91% were used. The reaction rate was essentially constant regardless of the relative humidity. I n all cases, enough water was present to react with the sulfur trioxide and to establish vapor pressure equilibrium with the drops without appreciably changing the initial relative humidity.
Droplets in Aerosol less than 1 Micron in Diameter
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When the concentration of sulfuric acid was sufficient, consecutive samples of the aerosol were taken with the impactor at several velocities less than sonic. From the amounts of acid collected, the mass percentage of aerosol greater than the characteristic diameter was calculated for each velocity. It was assumed t h a t the mass of aerosol less than 0.11 micron that was not collected at sonic velocity was negligible. The data gave a straight line when plotted on logarithmic probability paper,
For the case when the oxidation is the only reaction occurring, and when the initial concentration of sulfur dioxide is essentially constant, the general rate equation becomes dCa/d6 = k G
(1)
Figure 4 shows a plot of C, versus 6; since the percentage conversion per hour is constant, n must equal 1 and IC equals 0.0068 hr. -1 Under these conditions therefore the reaction is first order with respect t o sulfur dioxide. This is consistent with the theory that the rate controlling step is the effective collision of a n activated sulfur dioxide molecule with an oxygen molecule under conditions in which t.he concentration of the oxygen is constant (8).
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Since the reaction is first order with respect t o sulfur dioxide, the effect of a secondary reaction of the ozone produced with thesulfur dioxide molecule must be negligible under the conditions used.
Presence of Nitrogen Dioxide Has little Effect on Reaction Rate
A series of runs was carried out in which 1 or 2 p.p.m. of nitrogen dioxide were added along with 10 and 20 p.p.m. of sulfur dioxide. I n the dark, 20 p.p.m. sulfur dioxide and 1.0 p.p.m. of nitrogen dioxide produced no sulfuric acid. Also, no aerosol was formed by irradiation of 1.0 p.p.m. nitrogen dioxide and air in the absence of sulfur dioxide. The results of the runs in which mixtures of nitrogen dioxide and sulfur dioxide were irradiated are shown in Figure 5. The straight lines represent the average values obtained from Figure 3 and the vertical lines show the spread of data used t o determine these average values. The data fall within, or only slightly above the limit of the range of data
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REACTION TIME, MINUTES
Figure 4. Oxidation of sulfur dioxide in sun lamp irradiation
hence the size distribution of the particles may be represented by the log-probability distribution function (4). The mass median diameter, Dmmd,and the geometric standard deviation, uQ,were calculated for the various runs. Since ug was usually 1.1 t o 1.3, the aerosols produced by the photochemical reaction were rather homogeneous. The mass median diameter under the conditions used varied from 0.19 t o 0.45 micron. I n these results, only the amount of sulfuric acid was determined and no measurement of the total weight of the drops including the water could be made. However, water condenses rapidly on the acid drops and the acid concentration can be assumed t o be in equilibrium with the relative humidity at all times (8). This permits calculation of the mist loading which, in all cases, was less than 8 micrograms per liter. The number of par-
INDUSTRIAL AND ENGINEERING CHEMISTRY
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AIR POLLUTION
titles per cc. varied from 0.19 x 106 to 0.47 X lo6. The number of particles gives some indication of the probable history of the aerosol. From the results, it appears that the number of particles per cubic centimeter remains essentially constant after the initial nucleation and continued gaseous reaction results in the formation and absorption of sulfuric acid molecules and water vapor necessary t o maintain vapor pressure equilibrium.
phototubes has been reported by Luckiesh (18) and by Koller and Taylor (11). The phototube was mounted in a periscopetype brass housing attached t o a revolvable base and kept normal to the rays of the sun.
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liquid Phase Oxidation
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Two runs were made a t 20 p.p.m. to determine the extent of absorption and oxidation of sulfur dioxide in the sulfuric acid aerosol formed by the photochemical reaction. I n these runs the sulfur dioxide-air mixture was irradiated for 2 hours, a sample of :he aerosol was taken, and the irradiation was stopped. The mixture was again sampled after 1 and 2 hours. I n both cases, the sulfuric acid content decreased between 3 and 9%. From these results, it can be concluded that the rate of oxidation of sulfur dioxide in 30% sulfuric acid drops of 0.3-micron diameter in the absence of catalysts is negligible compared to the rate of photochemical oxidation. Liquid phase oxidation, however, may become important in water fogs, especially in the presence of dissolved catalysts.
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Figure 5.
Effect of nitrogen dioxide on reaction rate
In the present work, the mass concentration of the aerosol was in the order of 3 mg. per cubic meter of gas and the number concentration of the droplets was above 3.5 X 105 per cc. A typical natural water fog contains about 300 mg. per cubic meter with an average particle size of 20 microns, or only about 70 particles per cc. Under these conditions both the photochemical acid fog and the natural water fog would have approximately the same liquid surface area per cubic centimeter of air-i.e., about 1 X 10-3 sq. em. However, the natural water fog would contain 90 times as much mass of liquid. Thus, if the total amount of sulfur dioxide absorbed controls the rate of oxidation in the liquid phase a faster rate of oxidation would be expected for a water fog than was found in the photochemical fog. This possibility is being investigated.
Photochemical Reaction in Natural Sunlight A few runs were made in natural sunlight using a 64-cubic-foot reaction chamber lined with saran film and covered with Plexiglas windows. During these runs, the ultraviolet intensity of the sun in the region 2950 to 3400 A. was recorded continuously by means of a special 5% cadmium-95% magnesium alloy phototube obtained from the General Electric Lamp Research Department. The relative spectral sensitivity of cadmium-magnesium alloy May 1955
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Ultraviolet intensity in sunlight
The relative sensitivity of the phototube in the region from 2950 to 3200 A. is nearly the same as the relative absorption of sulfur dioxide a t room temperature. The phototube, therefore, gives a direct measure of the relative amount of ultraviolet light absorbed by sulfur dioxide. Thus, except for any variation of quantum yield between 2950 and 3200 A., the reaction rate should be proportional t o the recorded intensity. The variation of ultraviolet intensity from 9 A.M. to 5 P.M. for typical clear, hazy, and overcast days is shown in Figure 6. The phototube was also used t o measure the relative intensity of the sun lamp used for artificial sunlight. Use of special filters showed that 85% of the effective radiation from the sun lamp through the Plexiglas top consisted of the 3131 and 3022 mercury lines and 15% was due to the 2967 line. The data for natural sunlight indicate that essentially all the effective radiation was between 3000 and 3200 A. Measurements were made of the ultraviolet intensity of the sun lamp in the indoor reaction chamber by making traverses across the diameter of the cylinder a t the top and bottom of the chamber. The walls of the Lucite cylinder reflected the ultraviolet light and gave an irregular distribution of intensity in the chamber. However, an approximate average was calculated which showed that the intensity of the ultraviolet, in the region from 2950 to 3200 A., in the sun lamp chamber was approximately three times as great as the noon sunlight on a clear day and approximately five times as great as in the natural sunlight reaction chamber.
Reaction Rate in Natural Sunlight The results of test runs in natural sunlight were somewhat uncertain because of the great dilution of the initial sulfur dioxide and the presence of an oily aerosol which came from the wooden chamber when temperatures rose above 100" F. Based on data obtained during the first 30 minutes of reaction time and before the temperature had risen to 100' F., the reaction rate was in the order of 0.1% conversion per hour. This value may be compared with the value o.23yO per hour obtained by dividing the average rate in the sun lamp chamber, 0.68, by three to account for the greater ultraviolet intensity in the artificial illumination.
Conclusions The photochemical oxidation of sulfur dioxide in air a t concentrations of 5 t o 30 p.p.m. is a first order reaction with respect t o sulfur dioxide concentration. The reaction rate is slow, amount-
INDUSTRIAL AND ENGINEERING CHEMISTRY
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ENGINEERING, DESIGN, AND PROCESS DEVELOPMENT ing t o about 0.68% conversion per hour in artificial light and’O.l t o 0.2% per hour in intense natural sunlight. Nitrogen dioxide at concentrations of 5 to 20% of the sulfur dioxide concentration had no significant effect on the reaction rate. The presence of sodium chloride nuclei and variation of the relative humidity in the range from 30 t o 90% also had no effect on the reaction rate. The particle size of the sulfuric acid aerosol produced from the reaction was very small, varying from 0.2 to 0.4 micron in diameter. The sulfuric acid aerosol produced from the photochemical gas phase oxidation of sulfur dioxide apparently is not important in reducing the visibility under conditions ordinarily found in the atmosphere. Results of transmissometer tests on the effect of sulfur trioxide in decreasing visibility have been reported by the Stanford Research Institute (15). Accordingly, if the rate of conversion is 0.1% in bright sunlight, air containing 1.0 p.p.m. sulfur dioxide would have t o be exposed for approximately 100 hours of solar radiation t o reduce the visibility to one mile. Lower concentrations would have to be exposed to sunlight for a proportionately longer time. These statements refer only t o the photochemical gas phase oxidation and different results may be obtained by oxidation in the liquid phase, such as in the presence of a natural water fog.
Acknowledgment This work was a part of t h e program in the University of Illinois Engineering Experiment Station on a study of stack gases.
It was supported in part by the Smoke and Fumes Committee of the American Petroleum Institute as Contract S P 9 .
Literature Cited (1) Aitken, John, Proc. Roy. Soc. ( L o n d o n ) , 32, 183 (1911).
(2) Blacet, F. E., IND.ENG.CHEM.,44, 1339 (1952). (3) Blacet, F. E., private communication, 1950.
(4) Drinker, P., and Hatch, T., “Industrial Dust,” McGraw-Hill
Book Co., New York, 1936. (5) Ellis, C., and Wells, A. A., “The Chemical Action of Ultraviolet Rays,” p. 330, Reinhold, New York, 1941. Forsythe, W. E., “Measurement of Radiant Energy,” 1st ed., p. 77, McGraw-Hill Book Co., Kew York, 1937. Gerhard, E R., and Johnstone, H. F., Anal. Chem., in press. Gillespie, G. R., Ph.D. thesis, Universitv of Illinois. 1953. Hill, R. A,, Trans. Faraday SOC.,20, 107 (1924). International Critical Tables, Vol. V, p. 270, hIcGraw-Hiil Book Co., New York, 1928. Xoller, H., and Taylor, A. H., J . O p t . SOC.Amer., 25, 184 (1935). Luckiesh, M., Taylor, A. H., and Kerr, G. P., J . FrunkZinInst., 238. Pearse, R. W. B., and Gaydon, A . G., “The Identification of Molecular Spectra,” 2nd ed., p. 213, Chapman and Hall Ltd., London, 1950. Ranr, W. E., and Wong, J. B., IND.ENG.CHEM.,44, 1371 (1952). Stanford Research Institute, “The Smog Problem in Los Angeles County,” Second Interim Report, 1949. ENG.CHEM., Thomas, M. D., Ivie, J. O., and Fett, T. C., IND. ANAL.ED., 18, 383 (1946). RECEIVED for review September 2 5 , 1964.
ACCEPTED December 30, 1954.
Sources of Air Pollution literature JANET 6. MURK Barrett Division, Allied Chemical & Dye Corp., I River Road, Edgewater,
N. 1.
A
bibliography of bibliographies on air pollution has been prepared. It contains references providing an excellent starting point for the layman in his survey of this rapidly growing field. It also points up more obscure work to those better versed in the field. Legislation, institutional work, and symposia are cited. Samples of the many headings used for indexing air pollution literature are listed. Publications devoting space to air pollution and nomenclature problems inherent in the topic are also considered.
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HE literature of air pollution has become profuse. I t s rate of growth is increasing progressively as population and industry multiply and centralize. At first-an awareness of the harmful effects of dust exposure dates back t o approximately 75 A.D. when Pliny referred to it (14)-the emphasis was placed predominantly on the nuisance factor of smokes; t h e economic and physiological aspects did not achieve wide attention until this last score of years. The monetary value of stack effluents has been recognized and their recovery, or prevention of their ever becoming effluents, has been effected in many instances by industry acting on its own initiative. The essence of the present feeling toward air pollution control lies in this direction-that is, industry’s cleaning up of its own accord. There are no federal laws concerning smoke, dust, or fume tolerances. Thus far the legislative control issue has been relegated to local authorities because of variations of the problem in diverse topographic, climatic, and industrially developed areas. For these same reasons, antithetically, our common political divisions-city, county, and state-are often transcended by the situation thereby giving rise to the unique International Joint Commission of the United States and Canada. 916
I n t h e U. S. the problem has been met by and large on t h e municipal level with a distinct commission or t h e local department of health handling functional duties, such as drafting control ordinances, issuing permits for industrial equipment, and offering information. According to the American Municipal Association, only four U. S. cities of more than 25,000 population have not established some air pollution control (3). House Resolution 3555, submitted by Congressman Ray of Staten Island, proposes that the cost of treatment works for the abatement of air and stream pollution be amortized a t an accelerated rate for income tax purposes. This bill is currently in committee. Other legislation to provide the Department of Public Health, Education, and Welfare with funds to conduct i t s own investigations as well as to farm out research contracts is yet t o be considered by t h e Congress. T h e government Interdepartmental Committee on Air Pollution with representatives from the Atomic Energy Commission, the National Science Foundation, and the Departments of Agriculture, Commerce, Defense, Health, Education, and Welfare, and Interior is attempting t o coordinate federal research and t o give assistance t o t h e local committees. The recently dedicated Taft Sanitary
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