Isolation and Identification of Nitrosonium Hydrogen Sulfate (NOHS04) as a Photochemical Reaction Product in Air Containing Sulfur Dioxide and Nitrogen Dioxide William H. Schroeder Wastewater Technology Centre, Environmental Protection Service, Environment Canada, P.O. Box 5050, Burlington, Ont., Canada
Paul Uronel' Chemistry Department, University of Colorado, Boulder, Colo. 80302
~
During laboratory investigations intended to elucidate the kinetics and mechanism of the photo-oxidation of sulfur dioxide in air, the presence of a white solid material was repeatedly observed in the lower portions of the reaction flasks following photolysis of dilute mixtures of sulfur dioxide and nitrogen dioxide in purified air. Sufficient material was prepared for analytical investigations by irradiating higher concentrations of SOz-NOz-air mixtures. From the results of melting point determinations, microchemical analysis, reactant stoichiometry, infrared spectrophotometry, mass spectrometry, and chemical behavior, the photochemical reaction product was identified as nitrosonium hydrogen sulfate. This compound is also known t o be formed in the absence of light in the lead chamber process for the manufacture of sulfuric acid. The chemical reactions thought to be occurring in this process are examined. A possible reaction scheme for the formation of NOHS04 in polluted atmospheres is proposed, and the implications of these findings for atmospheric chemistry are discussed. H
Oxidation products of sulfur dioxide (sulfates and sulfuric acid droplets) constitute a significant fraction of the atmospheric aerosol over much of the world, both in the troposphere and in the stratosphere (1-4). The role of gaseous pollutants in the formation of the stratosphere aerosol layer has been discussed by Manson et al. ( 5 )and more recently by Bigg et al. (6). Stephens and Price ( 7 ) reported that aerosol generated photochemically appears to account for most of the visibility loss in smog as a result of light scattering by particles in the 0.3-l.O-~m-diameterrange. During the course of investigations into photochemical reactions of sulfur dioxide and nitrogen dioxide a t concentrations of 0.1-1% (v/v) in air (8),the presence of a white solid material was repeatedly observed to form in the gas phase and settle on the lower portions of the reaction flasks following irradiation of SOz-NOz-air mixtures. The formation of a solid reaction product by photolysis of SOz-NOZ-air mixtures was of considerable interest and pro~~
Present address, Environmental Engineering Sciences Department, University of Florida, Gainesville, Fla. 32611. 0013-936X/78/0912-0545$01 .OO/O
0 1978 American Chemical Society
vided valuable insight into concurrent studies dealing with the kinetics and mechanisms of thermal and photochemical reactions between sulfur dioxide and nitrogen dioxide in air (9,10).This paper describes the methods and procedures used to isolate and identify the white solid photochemical reaction product as nitrosonium hydrogen sulfate.
Experimental Reagents and Materials. The sulfur dioxide and nitrogen dioxide employed in these experiments were supplied in lecture bottles by the Matheson Co., Inc., and their purity was ascertained prior to use. All other chemicals were of reagent or analyzed reagent quality unless otherwise stated. The air used to prepare gaseous mixtures was obtained by passing ambient air through a series of traps containing anhydrous calcium sulfate (Drierite), indicating silica gel, molecular sieve (Type 5A), activated charcoal, anhydrous magnesium perchlorate (Anhydrone),and phosphorus pentoxide (Aquasorb). (Air purified in this manner still contains trace quantities of residual water vapor.) Laboratory supplied distilled water was redistilled in an all-glass Corning AG-3 still in the presence of potassium permanganate and distilled once again from a second all-glass apparatus. Apparatus and Procedure. The apparatus used for preparing gaseous mixtures of pollutants in air has been described previously (11). The gaseous mixtures contained in 2-L spherical borosilicate glass flasks were irradiated for controlled periods of time in a Srinivasan-Rayonet-Griffin photuchemical chamber reactor (The Southern New England Ultraviolet Co., Middleton, Conn.) having a bank of sixteen 24-W intermediate pressure ultraviolet lamps, coated to give a spectral distribution from 3100 to 4200 8, with an energy maximum a t 3500 A. The photon flux was measured by liquid-phase chemical actinometry using the potassium ferrioxalate system developed by Parker and Hatchard (12, 13) and following a procedure similar to that outlined by Calvert and Pitts (14). Note that an alternative and considerably less time-consuming approach to the preparation of the photolyte is described in the literature ( 1 5 ) . The photon flux delivered by the reactor lamps was found to be 1.9 X 10'6 photons s-l cm-2 and corresponded to about 10 times the intensity of natural sunlight a t midday in the lower atmosphere over this range of wavelengths (16). Volume 12, Number 5, May 1978
545
All infrared spectra were recorded either with a PerkinElmer 337 or a Beckman IR-12 spectrophotometer. Gas-phase infrared spectra were obtained with the aid of borosilicate glass gas sample cells having an effective path length of 10.0 cm and an inside diameter of 2.0 cm. Cell windows made of NaC1, KBr, or Irtran-2 (17,18) (available from Barnes Engineering Co.) were attached to the cell body with Apiezon-W wax. Mass spectral data were gathered using Varian-MAT CH-5 and CH-7 mass spectrometers. Preparation and Isolation of NOHSOd. Nitrosonium hydrogen sulfate was produced by irradiating mixtures of sulfur dioxide gas and nitrogen dioxide gas in an excess of purified air. The gaseous mixtures were prepared by introducing 200 mL of sulfur dioxide gas and 200 mL of nitrogen dioxide gas (volume corrected for the NzO4 equilibrium concentration), both at atmospheric pressure, into a 2-L spherical borosilicate glass flask containing purified air at atmospheric pressure. The flask was then irradiated until no further appreciable increase in the quantity of NOHS04 occurred. Nitrosonium hydrogen sulfate is a very hygroscopic material and must be handled and stored in the absence of moisture. The first attempt a t its isolation failed, because the product was exposed to room air. Almost immediately after exposure, the solid evolved a brown gas (nitrogen dioxide) and in a matter of minutes had turned into a colorless, viscous liquid. In later experiments care was taken to prevent moisture from coming into contact with the NOHS04. Isolation, storage, and manipulation took place in a glove bag filled with prepurified, dried nitrogen and containing a large dish of phosphorus pentoxide as a desiccant.
Results A variety of chemical, physical, and physicochemical methods of analysis were brought to bear upon the problem of identifying the white solid reaction product. Microchemical analysis for sulfur, nitrogen, and oxygen indicated a 1:1:5 stoichiometry for these three elements. In a separate experiment the molar ratio of sulfur to nitrogen was determined to be 0.97 f 0.05. Calculating the amount of oxygen present in the compound by difference again yielded a stoichiometric relationship of SN05. This is in agreement with the ratio of these three elements in NOHSOI, the chemical formula for nitrosonium hydrogen sulfate. A 1:l
0 SULFUR DIOXIDE
A NITROGEN DIOXIDE
z G
E
0.6.
z
5 0.5. z
0
0
0.4
I
I
I
I
I
I
2.0
4.0
6.0
8.0
10.0
12.0
IRRADIATION TIME (HOURS) Figure 1. Photochemicalreaction ratedata for 1.0% SO1 -k 1.0% NO2 in air 546
Environmental Science 8 Technology
WAVELENGTH (MICRONS) 15.0
20.0 25
1172 "I
3
E c
201
l
o
4000
3500
3000
2500
2000
1500
A V I 1200
,
1
,
1000
1
800
,
600
~
,
400
FREQUENCY (crn-'1
Figure 2. Infrared absorption spectrum of photochemical reaction product in KBr pellet
stoichiometric relationship between sulfur and nitrogen in the photochemical reaction product is also indicated on the basis of the observed decreases in concentration for sulfur dioxide and nitrogen dioxide in the gas phase during periods of irradiation (Figure 1). This shows that equal volumes of the reactants were consumed when gaseous mixtures consisting of 1.0%by volume of sulfur dioxide and nitrogen dioxide each in purified air were exposed to near ultraviolet radiation simulating the solar spectrum in the lower atmosphere. Melting point determinations were performed on the solid photochemical reaction product using a Mel-Temp Laboratory Devices melting point apparatus. The bulk of the sample was observed to undergo a phase change over the temperature range of 78-81 "C, but a gradual transition in appearance from a snow-white powder to a transparent, glassy-looking substance started to occur at 75 "C. For nitrosonium hydrogen sulfate, Moeller (19) reported melting with decomposition at 75 "C, Angus and Leckie (20)gave a range of 73-76 "C, and Tilden (21)a temperature range of 85-87 "C. A convenient laboratory procedure for its preparation was described by Biltz and Biltz (22).According to a number of sources (23-25), nitrosonium hydrogen sulfate is decomposed upon melting to give a colorless anhydride having the simplest molecular formula, SzNz09. The transition occurring during the melting process is shown in Equation 1: 2NOHSO4 .--+ (N0)2Sz07
+ H2O
(1)
The anhydride has a melting point of 217 "C and boils at 360 "C without decomposition. It is probably this anhydride that was observed during various experimental investigations of the thermal reactions of SO2 and NO, at elevated temperatures (21-28). Figure 2 is a reproduction of the infrared absorption spectrum obtained when a small quantity of the solid photochemical reaction product was mixed with infrared quality potassium bromide and pressed into pellet form. Because the infrared absorption spectrum for nitrosonium hydrogen sulfate does not appear to have been previously reported in the literature, a direct comparison of the spectrum recorded for the photochemical reaction product with that of NOHS04 was not possible. It was thought to be of interest, however, to compare the infrared absorption spectrum for the photochemical reaction product with that reported in the literature for potassium hydrogen sulfate, which might be expected to exhibit similar spectral characteristics. In Table I the major infrared absorption peaks observed for the photochemical reaction product are compared with those for KHS04 between 800 and 3000 cm-1 as published by Miller and Wilkins (29). The high degree of similarity for the IR spectral characteristics of the two substances is readily apparent from this table. Mass spectrometry was also used as a tool for the identification of the photochemical reaction product. Mass spectra obtained under two different analytical conditions are shown in Figure 3. For the first of these mass spectra (Figure 3A), a
Table I. Comparison of Frequencies and Relative Intensities of Infrared Absorption Bands for Photochemical Reaction Product (KBr Pellet) and Potassium Hydrogen Sulfate (29la Photochemical reaction product, Potassium hydrogen sulfate, 800-3000 cm-' 400-4000 cm-' 460 m 0-s-0 ... ... Deformation ... ... 579 5 ... ... 619 m 1 t
674 vs, s p
... ...
...
S-O(H)
852 s 875 s, sh 885 s
Stretch
c t
1010 s 1069 s, sp
Symmetric
SO; stretch
...
820 848
w, sh s
877
s
1005 1065
5 ' I s P
J -
t Asy r n m e t r i c SO3- stretch
1172 vs, b 1285 m
J -
1328 m , sh !
t
...
2 X (S-OH)? 4 t SOH 1 SOH
... ...
c
.. ...
. . 2911.. . 2951 . . ,
t SOH
341 1
H2O
vs, b
1325
w, sh
1640
w
S
2330 m , vb
2440
-
N0'
2482.. 2594..
1160 1280
-
J. -
...
..
2600
vw
2900
s, vb
... ...
a b = broad, m = medium, s = strong, s h = shoulder, sp = sharp, vb = very broad, vs = very strong, vw = very weak, w = weak.
(A) NO
60 -
w
0
1
iI
20
z
8
'0
1G
20 30 40 50 60 70 80 90 100
glass capillary containing the sample was broken inside the inlet port ofthe Varian-MAT CH-7 mass spectrometer. In this instance, the sample inlet port was at a temperature of 160 O C , which is high enough to convert nitrosonium hydrogen sulfate to its anhydride, (N0)2S207. The water of hydration thus liberated (see Eauation 1) would account for the Deak observed a t mle = 18. The base peak for the spectrum obtained under these conditions occurs a t a mass-to-charge ratio of 30, and in all likelihood is due to the species NO. The large peak a t mle = 28 is thought to result from the air (N2) coexisting with the sample in the glass capillary as well as residual air in the instrument. The presence of air would also contribute to the molecular oxygen (02) peak at mle = 32, which one could expect to be generated through fragmentation of the substance under consideration. The only sulfur-bearing entities detected were SO and SO2 a t mle = 48 and 64, respectively. For the scan of the other mass spectrum displayed (Figure 3B), the sample was contained in a gold cup, and the direct probe of the Varian-MAT CH-5 mass spectrometer was employed. In this case, except for peaks due to molecular nitrogen and oxygen (mle = 28 and 32, respectively), the only major peaks in the mass spectrum were those associated with sulfur containing entities and spanned the range from mle = 48 (SO) to mle = 98 (H2S04).The absence of a peak for NO (mle = 30) in this spectrum was due to the method of sample handling used in this particular mass spectrometric determination. It was not possible to keep the sample completely dry prior to the actual characterization. The sample, contained in an open gold cup, was inserted into the direct probe of the mass spectrometer, which was then evacuated to the normal operating pressure of the instrument (lov6torr) before the spectrum was scanned. Exposure of the sample in the gold cup to room air prior to insertion into the mass spectrometer caused the evolution of NO,, which was pumped off during evacuation. Nitrosonium hydrogen sulfate is hygroscopic and readily dissociates in the presence of water according to the following reactions (23, 30): H 2 0 t 2NOHS04 2H20
+ 3NOHS04
60
so
,
,
I
10
20
30 40 m/e
Figure 3. Mass spectra of solid photochemical reaction product from S02-N02-air mixtures ( A ) With sample inlet at 160 'C. (B) With direct probe at 90 'C
-
NO t NO2 t 2H2S04
"0.3
t 2 N 0 + 3H2S04
(2)
(3)
Combining the structural information contained in the two mass spectra results in a composite fragmentation pattern akin to that expected for NOHS04 and lends additional support to our contention that the photochemical reaction product is indeed nitrosonium hydrogen sulfate. In connection with other studies conducted in our laboratories a t that time (31),it was learned that nitrosyl chloride, NOCL, is formed by bringing into contact nitrosonium hydrogen sulfate and dry sodium chloride (32,33).The chemical reaction involved is: NOHS04
IT
-
+ NaCl
-
NaHS04
+ NOCl
(4)
This interesting reaction provided another opportunity to corroborate the identity of the white, solid reaction product formed when gaseous mixtures containing sulfur dioxide and NO2 purified air are exposed to near ultraviolet light simulating solar radiation. Thus, if this material is indeed nitrosonium hydrogen sulfate, its admixture with sodium chloride should generate gaseous nitrosyl chloride, which can be readily detected by its characteristic infrared absorption spectrum. The experimental verification proceeded as follows. An approximately equal quantity of dry powdered sodium chloride was added to a clean 2-L reaction flask containing some of the photochemical reaction product prepared previously, and a procedure similar to that described by Tilden (21) was followed. Figure 4 is a reproduction of the infrared absorption spectrum of a gas-phase sample withdrawn from the reaction flask after the flask and its contents were heated Volume 12, Number 5, May 1978
547
100
0 2
..-. 80-
w
0 2
60-
E
I
+
+
0
f
-.
+ hu NO + O(3P) O(3P) + + M O3 + M O3 + NO NO2 + O2 NO2
-
40
-
NOCI
03
(I)
z a
+ NO2
---*
+
l
l
1
1
NOCl
1
L
1
i
i
1
I
NOCl
1
2NO2
NO3
+0 2
+ NO2 + M N205+ M SO2 + hu (3400-2900 A) 'SO2 SO2 + hu (3900-3400 A) 3S02 'SO2 + so2 2s02 3S02+ SO2 NO3
f$ 200
+ 2N0
-
+
+
+
+
'SO2
+
SO2 + h u F
+
to about 80 "C for one-half hour. As expected, nitrosyl chloride was the major gaseous product in the reaction flask. The small quantity of nitrogen dioxide that was also present (Figure 4) was probably the result of photodecomposition of some of the NaOCl and subsequent oxidation of the liberated nitric oxide. Nitrosyl chloride is photodissociated at wavelengths below about 7600 A ( 1 4 ) .
Discussion Nitrosonium hydrogen sulfate appears to have been known has long been aca t least since 1755 (34).Whereas "SO5 cepted as the empirical formula for nitrosonium hydrogen sulfate, it was not until 1926 that Elliot et al. (25)convincingly showed its composition to be HOS020NO. Reliable evidence indicating an ionic structure for this compound has been obtained by a number of investigators using a variety of techniques such as electrolysis (eo),conductivity studies (35,36), cryoscopic measurements (35-38), Raman spectroscopy (20, 39-41 ), and physicochemical considerations (42). Furthermore, a considerable amount of work has been undertaken to prove the existence of nitrosonium hydrogen sulfate as an intermediate species in the chemical reactions involved in the manufacture of sulfuric acid by the traditional lead chamber process (43). The most essential reactions believed to be occurring in the lead chamber process for the manufacture of sulfuric acid are:
-
+ SO2 + H20 NO + Has04 2H2S04+ NO + NOz 2NOHS04 + H 2 0 2S02 + 3N02 + H2O 2NOHS04 t NO 2S02 4-NO + NO2 + 0 2 + H20 2NOHSOd 2NOHSO4 + H20 2H2S04 + NO + NO2 NO2
+
---*
+
(5) (6)
(7)
(8) (9)
These reactions only indicate the molecular entities involved in each step, and do not allow any conclusions to be drawn about the actual mechanistic details. In the presence of a sufficient quantity of water, Reactions 5 , 6, and 9 predominate, leading to the formation of sulfuric acid. In this particular sequence of reactions, the nitrogen oxides (NO and NO2) participate in the reaction scheme without being consumed, thus fulfilling, by definition, the role of a catalyst. With an insufficient quantity of water present (representing a situation in which H20 is the limiting reagent), Reactions 7 and 8 play the dominant roie, thus resulting in the accumulation of NOHS04 in the lead chamber (43). In addition to the chemical reactions just described, certain photochemical processes are also expected to be occurring under our experimental conditions (44-48). These include: 548
Environmental Science & Technology
so2
'SO2 + 3S02 3S02 SO2 + huF
- so2 + +
3S02 SO2
2S02
SO3 + SO
so3 + H20
-
H2S04
In considering the possibilities for chemical interaction between substances formed by the photolysis of NO2 (i.e., 0, 03,NO, and NO3) and sulfur dioxide (in its ground state), note that an estimate for the upper limit of the rate of the reaction: NO3 + SO2
-
NO2
+ SO3
(22)
has recently been determined by Daubendiek and Calvert (49).On the basis of their findings they suggested that Reaction 22 does not constitute an important removal path for SO2 in polluted atmospheres. Previous work by Wilson et al. (50) had indicated that NO3 or N205 could rapidly oxidize sulfur dioxide. The reaction between sulfur dioxide and atomic oxygen:
SO2
+0+M
-*
SO3 + M
(23)
has been studied by Mulcahy et al. (51).They found that the rate constant for the reaction between sulfur dioxide and atomic oxygen is smaller than the rate constant for the reaction of 0 atoms with NO2 and NO. However, by irradiating gaseous mixtures containing 2 ppm of SO2 and nitrogen dioxide at concentrations in the range of 0-10.2 ppm in air, Smith and Urone (52) showed that SO2 can compete effectively for the atomic oxygen with other species in this system so long as the S02/N02 concentration ratio is greater than unity. Thus, sulfur dioxide at 2 ppm in air in the presence of 0.85 and 1.7 ppm NO2 was photo-oxidized at a rate of about four times that measured for SO2 in air alone. But at NO2 concentrations of 3.4, 5.1, and 10.2 ppm, the rate of photooxidation of SO2 was similar to that observed for SO2-air mixtures in the absence of nitrogen dioxide. Sulfur dioxide in its electronic ground stage does not react at an appreciable rate with either O2 or 03.Similarly, any reaction between nitric oxide and sulfur dioxide should be insignificant in relation to other possible pathways for the conversion (oxidation) of sulfur dioxide in irradiated air mixtures containing nitrogen oxides and sulfur dioxide. It is not yet entirely clear what role is being played by photochemical activation processes and photolytically generated reaction intermediates in enhancing the rate of con-
version of sulfur dioxide to NOHS04 in irradiated systems over that observed in the absence of light (8). A possible mechanism which could contribute to, or even entirely account for, the formation of nitrosonium hydrogen sulfate in our irradiated mixtures is given below: 2N02 NO
+ H2O
+
-
HONO2
+ NO2 + H2O
-
+ HONO
2HONO
+ HzO 2HON02 + NO HON02 + hv HO + NO2 HONO + hu HO + NO HO + SO2 HOSO2 HOSO2 + NO2 HOS020NO
3N02
+
--*
+
+
(25) (26)
(27)
(28) (29)
(30)
To the best of our knowledge, these investigations (8) constitute the first time that nitrosonium hydrogen sulfate has been recognized, isolated, and conclusively identified as a reaction product from the photolysis of sulfur dioxide-nitrogen dioxide-air mixtures. In a recent study of the N205SO2-O3 reaction system, Daubendiek and Calvert (49) observed that gaseous SO3 and NO2 react rapidly to form a relatively nonvolatile white solid and speculated that it may be the same as the white solid photochemical reaction product described by Urone et al. (IO) in a more cdmplex system involving irradiation of dilute gaseous mixtures of SO2 and NO2 in purified air. Because of the absence of air, moisture, and solar radiation in the experiments performed by Daubendiek and Calvert, it is likely that their nonvolatile white solid is a member of the pyrosulfate family of compounds, viz., (NOX)2S207.Various members of this family of compounds, consisting of sulfur, nitrogen, and oxygen, have been observed by different investigators (26-28, 53-55) during laboratory studies of the thermal reactions between sulfur oxides (SOz,SO3) and nitrogen oxides (NO, N02, Nz03, N205). Under the conditions of our experiments, involving irradiation of sulfur dioxide-nitrogen dioxide-air mixtures, nitrosonium hydrogen sulfate was the principal reaction product. In ambient air, where the concentration of water vapor is generally several orders of magnitude greater than the concentration of individual air pollutants, any NOHS04 formed as a result of the photochemical reaction between sulfur dioxide and nitrogen dioxide is expected to be hydrolyzed immediately with the resultant formation of sulfuric acid aerosol. As a transient intermediate, NOHS04 would be extremely difficult, if not impossible, to detect in the atmosphere. It remains to be pointed out that, because our laboratory investigations were conducted at concentrations several orders of magnitude greater than normally encountered in polluted urban atmospheres, the chemical and photochemical reactions postulated for the formation of NOHS04 may not necessarily represent the predominant mechanism for the photo-oxidation of SO:! in the atmosphere. Nevertheless, Calvert and McQuigg (56) have proposed a similar reaction scheme involving the formation and hydrolysis of NOHSO4 as well as N02HS04 to account for the conversion of sulfur dioxide to aerosol in polluted urban atmospheres.
Acknowledgment The authors express their appreciation to A. Lazrus of the National Center for Atmospheric Research, Boulder, Colo., for the interest shown in these investigations and for providing us with the results of a microchemical analysis for a sample of nitrosonium hydrogen sulfate generated in our laboratories.
Literature Cited (1) , , “Chemical Reactions in Urban Atmosoheres”. C. S. Tuesdav. Ed.. Elsevier, New York, N.Y., 1971. (2) Urone. P.. Schroeder. W. H.. Enuiron. Sci. Technol.. 3. 436 (1969). (3) Bufalini, M., ibid., 5,685 (1971). (4) Cadle, R. D., in “Aerosols and Atmospheric Chemistry”, G. M. Hidy, Ed., Academic Press, New York, N.Y., 1972. (5) Manson, J. E., Junge, C. E., Chagnon, C. W., in “Chemical Re-
I
I
actions in the Lower and Upper Atmosphere”, Interscience, New York, N.Y., 1961. (6) Bigg, E. K., Ono, A., Thompson, W. J., Tellus, 22,550 (1970). ( 7 ) Stephens, E. R., Price, M. A., in “Aerosols and Atmospheric Chemistry”, G. M. Hidy, Ed., Academic Press, New York, N.Y., 1972. (8) Schroeder, W. H., PhD thesis, University of Colorado, Boulder, Colo., 1971. (9) Schroeder, W. H., Urone, P., paper presented before the Div. of
Water, Air and Waste Chemistry, American Chemical Society, 160th National Meeting, Chicago, Ill., Sept. 13-18, 1970. (10) Urone, P., Schroeder, W. H., Miller, S. P., Proceedings of the Second International Clean Air Congress, Washington, D.C., Dec. 6-11,1970. (11) Urone, P., Lutsep, H., Noyes, C. M., Parcher, J. F., Enuiron. Sci. Technol., 2,611 (1968). (12) Parker, C. A., Proc. Roy. SOC. (London),A220,104 (1953). (13) Hatchard, C. G., Parker, C. A., ibid., A235,518 (1956). (14) Calvert, J. G., Pitts, J. N., Jr., “Photochemistry”, Wiley, New York, N.Y., 1966. (15) Baxendale, J. H., Bridge, N. K., J. Phys. Chem., 59, 783 (1955). (16) Leighton, P. A., “Photochemistry of Air Pollution”, Academic Press, New York, N.Y., 1961.
(17) Colthua N. B.. Dalv. L. H.. Wiberlev. S. E.. “Introduction t o Infrared a i d Raman Spectroscopy”, Academic Press, New York, N.Y., 1964. (18) Bent, R., Ladner, W. R., Pankhurst, K. S.,Waite, B. D., Nature, 193,62 (1962). (19) Moeller, T., J . Chem. Educ., 23, 441, 542 (1946); 24, 149 (1947). (20) Angus, W. R., Leckie, A. H., Proc. Roy. SOC. (London), A149, 327 (1935); Nature, 134,572 (1934); Trans. Faraday Soc., 31,958 (1935). (21) Tilden, W. A., J . Chem. SOC.,27,630 (1874). (22) Biltz, H., Biltz, W., “Laboratory Methods of Inorganic Chemistry’’, Wiley, New York, N.Y., 1928. (23) “Gmelins Handbuch der Anorganischen Chemie: Schwefel, Teil B, System No. 9”, Verlag Chemie GmbH, WeinheimDergstr., West Germany, 1953. (24) Michaelis, A,, Schumann, O., Ber., 7, 1075 (1874). (25) Elliot, G. A., Kleist, L. L., Wilkins, F. J., Webb, H. W., J . Chem. SOC.,1926, p 1219. (26) Bruening, A,, Liebigs Ann. 98,377 (1856). (27) Terres, E., Constantinescu, M., 2. Angem. Chem., 47, 470 (1934). (28) Boreskov, G. K., Illarionov, V. V., Z. F i t . Khim., 14, 1428 (1940). (29) Miller, F. A,, Wilkins, C. H., Anal. Chem., 24,1253 (1952). (30) Yost, D. M., Russell, H., Jr., “Systematic Inorganic Chemistry”, pp 48-9, Prentice-Hall, New York, N.Y., 1944. (31) Schroeder, W. H., Urone, P., Enuiron. Sci. Technol., 8, 756 (1974). (32) Coleman, G. H., Lillis, G. A., Goheen, G. E., in “Inorganic Syntheses, Vol I”, p p 55-9, McGraw-Hill, New York, N.Y., 1939. (33) Burns, W. G., Bernstein, H. J., J. Chem. Phys., 18, 1669 (1950). (34) Bernhardt, J. C., “Chymische Versuche und Erfahrungen”, p 129, Leipzig, 1755. (35) Hantzsch, A., Berger, K., 2. Anorg. Chem., 190,321 (1930). (36) Kortnev, A. V., Bol’shakov, A. G., Aleksenko, N. A., Gasyuk, G. N.. Naucnve ZaDiski Odessk. oolitechn. Inst.. 20.3 (1960): Chem. Adstr., 19438 (f961). (37) Hantzsch. A.. 2. Phvs. Chem.. 65.41 (1909). (38) Gillespie, R. J., Graham, J., Hughes, E. D., Ingold, C. K., Peeling, E.R.A., J . Chem. Soc., 1950, p 2504. (39) Chedin, J., Compt. Rend., 200, 1397 (1935). (40) Millen, D. J., J . Chem. SOC., 1950, p 2600. (41) Simon, A., Richter, H., J . Prakt. Chem., 5,68 (1958). (42) Seel, F., Z. Anorg. Chem., 261,75 (1950); Angetu. Chem., 68,272 (1956). (43) Ridler, E. S., in “The Manufacture of Sulfuric Acid”, W. W. Duecker and J. R. West, Eds., Reinhold, New York, N.Y., 1959. (44) Altshuller, A. P., Bufalini, J. J., Photochem. Photobiol., 4, 97 (1965). I
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Volume 12, Number 5, May 1978 549
(45) Altshuller, A. P., Bufalini, J. J., Enuiron Sci. Technol., 5 , 39 (1971). (46) Finlayson-Pitts, B. J., Pitts, J. N., Jr., in “Advances in Environmental Science and Technology”, J. N. Pitts, Jr., and R. L. Metcalf, Eds., Vol7, Wiley, New York, N.Y., 1977. (47) Collier, S. S., Morikawa, A., Slater, D. H., Calvert, J. G., Reinhardt, G., Damon, E., J . Am. Chem. SOC.,92,217 (1970). (48) Sidebottom, H., Badcock, C., Jackson, G. E., Calvert, J. G., Reinhardt, G. W., Damon, E. K., Enuiron. Sci. Technol., 6, 72 (1972). (49) Daubendiek, R. L., Calvert, J. G., Enuiron. Lett., 8, 103 (1975). (50) Wilson, W. E., Jr., Levy, A., Wimmer, D. B., J . Air Pollut. Con-
trol Assoc., 22, 27 (1970). (51) Mulcahy, M. F., Stevens, J. R., Ward, J. C., J . Phys. Chem., 71, 2124 (1967). (52) Smith, J. P., Urone, P., Enuiron. Sci. Technol., 8,742 (1974). (53) Stopperka, K., Kilz, F., 2. Chem., 8,435 (1968). (54) Stopperka, K., Wolf, F., Suess, G., 2. Anorg. Allg. Chem., 359, 14 (1968). ( 5 5 ) VandorDe, B., Heubel, J.. C o m ~ tRend.. . 260.6619 (1965). (56) Calvert; J. G., McQuigg, R. D.;J. Chem. Kinet., Symp. No. 1, Suppl. to Vol7,113-54 (1975). Received for review March 8, 1977. Accepted October 25,1977.
Effects of Continuous H2S Fumigation on Crop and Forest Plants C. Ray Thompson* and Gerrit Kats University of California, Statewide Air Pollution Research Center, Riverside, Calif. 9252 1
Continuous fumigation of alfalfa (Medicago satiua L ) , Thompson seedless grapes (Vitis vinifera), lettuce (Lactuca sativa), sugar beets (Beta uulgaris), California buckeye (Aesculus californica), ponderosa pine (Pinus ponderosa), and Douglas fir (Pseudotsuga menziesii) with 3000 parts per billion (ppb) H2S in greenhouses caused leaf lesions, defoliation, reduced growth, and death of sensitive species. Three hundred ppb caused lesser but similar effects. Sulfur accumulated in leaves depending upon dosage. Faster growing plants accumulated sulfur more rapidly. Lower levels of HZS, 30 ppb and sometimes 100 ppb, caused significant stimulation in growth of lettuce, sugar beets, and alfalfa. The stimulation occurred at certain times of year and may be influenced by temperature and/or humidity. Hydrogen sulfide is a malodorous gas emitted from many industrial and biological processes. It is highly toxic to man ( I ) , but its effects on vegetation have received scant attention. Two high concentration short-term fumigation studies represent the major works done with this compound. McCallan et al. ( 2 ) fumigated 29 species of vegetation in greenhouses with 20-400 ppm H2S for 5 h during the middle of the day. A wide difference in degree of injury was observed. No effect was seen on eight species at 400 ppm, while some plants were injured by less than 40 ppm. Young tissue was most sensitive. Increasing temperature caused a rapid increase in injury. Benedict and Breen (3) fumigated 10 common weed species with a number of air pollutants including H2S. They used 100-500 ppm for 4 h, and the fumigations were done on plants 3 and 6 weeks of age. They observed that the younger plants were more susceptible and that drier soil caused plants to show much more injury. These short-term studies used H2S concentrations at least one order of magnitude higher than other compounds such as S02, fluorides, and ozone which are normally encountered in polluted air; yet, the degrees of injury to vegetation were comparable. As a result, it has been assumed by investigators that the amounts which occur near industrial emitters or from natural sources cause minor impacts. Thus, oxidized sulfur, as SOz, is much more toxic at high concentrations on a molar basis than HzS, but at low levels it can overcome sulfur deficiency and serve as a fertilizer (4-6). No controlled studies have been done to find out what the effects of long-term, ambient levels of H2S may be on vegetation. Because development of geothermal energy can cause emissions of considerable amounts of HZS, the present study was designed to assess its effects on several crop and forest species which grow near sources of geothermal energy. 550
Environmental Science & Technology
Procedure Fumigations. Continuous, uniform fumigations were conducted in four greenhouses glazed with translucent corrugated fiberglass (Glassteel, Duarte, Calif.). All greenhouses were equipped with activated charcoal filters on air intake and exhaust to exclude ambient pollutants and avoid cross contamination. Temperatures were maintained within 3 “C of ambient during hot weather with evaporative coolers. During cool weather, inside and outside temperatures were the same. Hydrogen sulfide diluted with nitrogen was injected into the carbon-filtered, incoming air stream to provide the required concentrations of fumigant. H2S was monitored with a Phillips Model 1900 H2S analyzer. Two fumigation procedures were employed. In the first series of experiments during 1975, alfalfa (Medicago sativa L ) , Thompson seedless grapes (Vitis vinifera), ponderosa pine (Pinus ponderosa), and California buckeye (Aesculus californica) were exposed to concentrations of 0,30,300, and 3000 parts per billion (ppb) of H2S. Lower pollution levels (0, 30,100, and 300 ppb) were used in the second fumigation series during 1976 which involved alfalfa, lettuce (Lactuca satiua), sugar beets (Beta vulgaris), Douglas fir (Pseudotsuga menziesii), and a second set of Thompson seedless grapes. Materials. All species of plants were grown in 28-L pots except lettuce and Douglas fir which were grown in 20- and 7-L pots, respectively. A soil mix consisting of peat moss, redwood shavings, and silt (1-1-1) was used to which salts were added at the following rates, in kg/m3 of mix: single super phosphate, 1.5; KNOB,0.15; K2S04, 0.15; dolomitic limestone, 2.2; and oyster shell lime, 0.90. Copper, Zn, Mn, and Fe were added at 30,30,15,and 15 ppm, respectively. The plants were irrigated with one-half strength Hoagland’s solution twice weekly. Tap water was used when additional irrigations were needed. Plant species, varieties used, and duration of treatments are summarized in Table I. All data, where replicated, were analyzed statistically by an analysis of variance and a multiple-range test (7). Results Effects of Continuous HzS Fumigation on Alfalfa. During 1975 when fumigation levels were 0,30,300,and 3000 ppb HzS, green alfalfa showed white marginal leaf lesions on mature leaves within 5 days at the highest level of H2S. To determine yield, the plants were cut at 28-35-day intervals depending upon growth rate and development. Growth was reduced during the first growing period at both 300 and 3000 ppb H2S with Hayden (Table 11). The highest level reduced growth of Eldorado, and a trend was shown at 300 ppb. No effect on growth was caused by 30 ppb. During the subsequent
0013-936X/78/0912-0550$01.00/0
0 1978 American Chemical Society
