curred subsequently to discharge to the lake, since mirexcontaminated sediments did not have a significant amount of 8-monohydromirex compared to mirex (2, 6 ) . Because a large proportion (>30%) of the mirex-related compounds in biota are photodegradation products, most of the mirex which entered the food web must have been held in the water column and recycled to the surface of the lake over a period of years. Alternatively, the rate of photolysis may have been enhanced by some unknown mechanism in association with dissolved organic matter, particle surfaces, or phytoplankton. The constancy of the ratio of 8-monohydromirex to mirex levels in Lake Ontario herring gul1,eggs over the years 1972 to 1978 is shown in Table IV. The ratio h samples of five species of fish (coho salmon, alewives, smelt, carp, and eel) from Lake Ontario in 1977 was also in the range 0.3-0.4 (6, 7). These data suggest that an equilibrium between input to the lake and photodegradation existed. However, the declines in mirex levels in herring gull eggs since 1974 ( 1 9 ) ,presumably associated with cessation of direct discharge to the lake, should have resulted in an increase in the ratio if photodegradation had been occurring throughout this period. That portion of the mirex and photomirex in the lake which is available for bioaccumulation in fish and herring gulls must therefore be protected from further photodegradation by being sequestered and cycled between and within various food webs over a period of several years. The data also indicates that mirex in the sediments is not being recycled into the ecosystem at an appreciable rate.
Sampson, R., J . Agric. Food Chem., 24,1189-93 (1976). (5) Hallett, D. J., Norstrom, R. J., Onuska, F. I., Comba, M. E., “Fate of Pesticides in Large Animals”, Ivie, G. W., Dorough, H., Eds., Academic Press. New York. 1977. uu 183-92. (6) Norstrom, R. J., Hallett, D. J.,’Soktegard, R. A., J . Fish. Res. Board Can., 35,1401-9 (1978). (7) . , Norstrom. R. J.. Won. H. T.. Holdrinet. M. V. H.. Calwav. P. G.. Naftel, C. D., J . Assoc. O f f . Anal. Chem.; 63,37-42 (1980j.’ (8) Carlson, D. A., Konyha, K. D., Wheeler, W. B., Marshall, G. P., Zaylskie, R. G., Science, 194,939-41 (1976). (9) Onuska, F. I., Comba, M. E., J . Chromatogr., 26, 133-45 (1976). (10) Dilling, W. L., Dilling, M. L., Tetrahedron, 23, 1225-33 (1967). (11) Alley, E. G., Layton, R., in “Mass Spectrometry and NMR
Spectroscopy in Pesticide Chemistry”, Hague, R., Biros, F. J., Eds., Plenum Press, New York, 1974, pp 81-91. (12) Hallett, D. J., Khera, K. S., Stoltz, D. R., Chu, I., Villeneuve, D. C., Trivett, G., J . Agric. Food Chem., 26,388-91 (1978). (13) Lane, R. H., Grodner, R. M., Graves, J. L., J. Agric. Food Chem.,
Literature Cited
24,192-3 (1976). (14) Alley, E. G., Layton, B. R., Minyard, J. P., J. Agric. Food Chem., 22,442-5 (1974). (15) Alley, E. G., Layton, B. R., Minyard, J. P., J. Agric. Food Chem., 22,727-9 (1974). (16) Alley, E. G., Dollar, D. A., Layton, B. R., Minyard, J. P., J. Agric. Food Chem., 21,138-9 (1973). (17) Holmstead, R. L., J . Agric. Food Chem., 24,620-4 (1976). (18) Ungefug, G. A., Scherer, K. V., Tetrahedron Lett., 33,2923-6 (1970). (19) Weseloh, D. V., Mineau, P., Hallett, D. J., Trans. North Am. Wildl. Nat. Resour. Conf., 44th 543-57 (1979). (20) Gilman, A. P., Fox, G. A., Peakall, D. B., Teeple, S. M., Carroll, T. R., Haymes, G. T., J. Wildl. Manage., 4,458-68 (1977). (21) Moore, F. R., Bird Banding, 47,141-59 (1976). (22) Waters, E. M., Huff, J . E., Gerstner, H. B., Enuiron. Res., 14, 212-22 (1977). (23) Gibson, J. R., Ivie, G. W., Dorough, H. W., J . Agric. Food Chem., 20, 1246-8 (1972). (24) Andrade, P., Jr., Wheeler, W. B., Carlson, D. A., Bull. Enuiron. Contam. Toxicol., 14,473-9 (1975). (25) Stein. V. B.. Pittman. K. A.. Bull. Enuiron. Contam. Toxicol.. 18,425-7 (1977). (26) Cripe, C. R., Livingston, R. J., Arch. Enuiron. Contam. Toxicol., 5, 295-303 (1977).
(1) Kaiser, K. L. E., Science, 185,523-5 (1974). (2) Holdrinet, M. V. H., Frank, R., Thomas, R. L., Hetling, L. J., J . Great Lakes Res., 4,69-74 (1978). (3) Kaiser, K. L. E., Enuiron. Sci. Technol., 12,520-8 (1978). (4) Hallett, D. J., Norstrom, R. J., Onuska, R. I., Comba, M. E.,
Received for review September 4, 1979. Accepted March 31, 1980. Part of the material in this study was presented at the International Symposium on the Analysis of Hydrocarbons and Halogenated Hydrocarbons, May 23-25,1978, Hamilton, Ontario, Canada.
Acknowledgments The authors wish to thank Dr. A. P. Gilman, in particular, for the complicated task of organizing the collection of the eggs from so many different sites, and Dr. B. R. Layton for the gift of hydromirex standards. M. J. Mulvihill is thanked for the sample preparation and analysis.
Aqueous Scrubber Reactions of Pyrrhotite (FeS) with Sulfur Dioxide George D. Case’”, Derry A. Green2, and Gerald W. Stewart3 Analytical and Supporting Research Division, Morgantown Energy Technology Center, U S . Department of Energy, Morgantown, W.Va. 26505
The aerobic hydrolysis of iron sulfides has long been a subject of environmental concern. Its importance in the chemistry of acid mine drainage ( I ) and ore smelting operations ( 2 ) has been documented. Thermodynamic properties of iron sulfides (3-6) and the kinetics of their aqueous reactions (7,8) have also been reported under carefully controlled laboratory conditions. However, the roles of iron sulfides in emerging advanced environmental control technologies have only recently been examined for their applicability to coal combustion (9-13) and low-Btu gasification of coal (14-20). Recently, FeS has been cited (9-13) as a possible substitute Present address, Resource Technologies Group, Inc., Route 2, Box 93A, Morgantown, W. Va. 26505. Present address, Atlantic Research Corp., 5390 Cherokee Ave., Alexandria, Va. 22314. Present address, Aerodyne Research, Inc., Bedford Research Park, Crosby Drive, Bedford, Mass. 01730. 866
Environmental Science & Technology
for limestone sorbent in flue gas SO2 removal from conventional coal-fired boilers. The present work investigates the actual chemistry of FeS reactions with SO2 under near-neutral pH conditions, and details the mechanisms by which SO2 is removed from the gas phase in aqueous scrubbers to produce soluble Fe2+ and SO4*-, as well as precipitated FeO(0H) as the only significant products (21). SO2 capture efficiency is found to depend on solution pH and buffer capacity, and is not affected by FeS.
Experimental All chemicals were reagent grade. Physical chemical analyses of the FeS starting material (21) showed it to consist primarily of hexagonal pyrrhotite of stoichiometric composition. Aerobic hydrolysis experiments to determine reaction products and pathways were carried out in open vessels as
This article not subject to U.S. Copyright. Published 1979 American Chemical Society
The use of pyrrhotite (FeS) in aqueous scrubber systems for SO2 control was examined for (a) its ability to remove SO2 from simulated flue gases and (b) the reaction mechanisms and ensuing prosducts. Principal products observed for FeS hydrolysis, with or without an SO2 source, include lepidocrocite (FeO(OH:I),soluble Fe2+,and soluble Sod2-. Extension of known thermodynamic functions and kinetic rate expressions to the present conditions give reasonable qualitative and quantitative fits to the observed data. No evidence is found for pyrite (FeS2) or elemental sulfur (SO) in the end products
of FeS-SO2 scrubber reactions, nor are these predicted from thermodynamic and/or kinetic considerations for the conditions studied. Measurements of SO2 uptake from gas streams by aqueous scrubbers, with or without FeS, indicate that 90% or better SO2 removal efficiency is achieved a t pH >5. Under these conditions, we could not establish any direct role for FeS in the removal of SOz. Any reaction between SO2 and the hydrolysis products is rate limited by the hydrolysis and is a t least 103-106 times slower than the space velocity of the gas through the scrubber a t 90% SO2 removal.
described previously (21). SO2 capture experiments were carried out under a controlled atmosphere in an air-tight titration apparatus, by passing gas containing SO2 through an aqueous suspension of FeS, followed by examination of the particles and analysis of the solutions. Total iron and ferrous iron (Fez+)were determined spectrophotometrically by the 1,lO-phenanthrolinemethod (22). Soluble sulfate (S042-) was determined tur bidometrically (23, 24), and thiosulfate (S20:j2-) was determined spectrophotometrically according to the method of IEriksen and Lind (25).Physical and chemical analyses of the insoluble reaction products were carried out by means of X-ray diffraction (XRD), scanning electron microprobe analysis (SEMI, X-ray fluorescence (XRF), and infrared spectrometry (IR) (26) as described previously (21). Reactions of SO2 gas with FeS were followed potentiometrically for changes in p H and oxidation-reduction (redox) potential (21, 27). Gas-phase SO2 was continuously monitored in situ by means of a Lear Siegler Model SM-400 second derivative spectrometer (24) using a l-cm path flow cell.
0.52:l.OO as opposed to 1.OO:l.OO in the FeS starting material. Atom ratios for S:Fe in the precipitate fractions are lower still (0.18:l.OOto 0.37:l.OO) but somewhat variable because of the proportion of finely divided, unreacted FeS present in the latter instance. The correlation is qualitatively demonstrated by scanning electron microprobe-XRF examinations of these preparations (21). Because most of the detected X-ray fluorescence intensity originates within the topmost 0.5-5.0 p m of the surface (28-30), present observations strongly suggest that unreacted residue particles are S deficient near the surface. Consequently, FeSz or Soformation by means of a surface reaction of FeS with SO2 or 0 2 is unlikely. Elemental analyses were also carried out on individual particles in the precipitate fractions by means of scanning electron microprobe-XRF analysis (21). Scanning electron micrographs of representative fields of view were compared to corresponding electron backscatter images and iron XRF and sulfur XRF element maps. Particles of iron sulfides were identified from their brilliant electron-electron backscatter images and were quantitatively analyzed for Fe and S content. All of the sulfur-rich regions in the S maps coincide with the brilliant surfaces. According to the iron maps, none of the sulfur-rich regions are deficient in iron. As seen by XRF microprobe analysis, the individual sulfur-rich particles all contain S:Fe atom ratios between 0.8:l and 1.1:l. Hence, all experiments suggest that these particles consist of FeS which may be S depleted with respect to pyrrhotite. No sulfur-rich regions other than FeS crystals have been found. Figure 1 presents infrared (IR) spectra of the insoluble reaction products of FeS hydrolysis under a variety of conditions. The presence of FeO(0H) as lepidocrocite is apparent in all of the precipitates as shown in spectra A-D (spectrum F is a lepidocrocite standard (21,31)).IR bands due to FeO hydrates (31,32)are observed in the hydrolysis products when either NH3 (C) or SOP (D) is present in the reaction medium. IR bands due to Fe2(S04)3.9H20 as coquimbite (31,32) are observed in the precipitate formed by FeS hydrolysis in the presence of bisulfites (spectrum D). This observation is consistent with the accumulation of dissolved iron sulfates such as FeS04+ and Fe(S04)z- during SO:! scrubbing with FeS (9). In no instances are IR bands due to FeS2 or Soobserved. Figure 1 confirms FeS as the major constituent of the unreacted residue fraction from the hydrolysis reaction; however, the weak bands at 3400 and 530 cm-l in spectrum E also point to an iron oxide hydrate as a significant minor product. Since the iron oxide bands are not present in IR spectra of the FeS starting material, it is reasonable to conclude that the iron oxide hydrate in the residue fraction is a hydrolysis product bound to the FeS surfaces. The X-ray fluorescence data, which show the surface to be iron rich and S depleted, confirm this identification. It is believed that the presence of an iron oxide layer on the FeS surface following aerobic hydrolysis (&SO,) may explain the fool's gold appearance of the residue fraction. There appears to be no evidence of significant pyrite (FeS2) or elemental sulfur formation under the present experimental conditions. Table I summarizes the material balances for iron and
Results Products from Iron and Sulfur in FeS Sorbent. One objective of this study was to investigate the role of FeS in converting flue gas SO2 to elemental sulfur. Previous work (2) has shown this pyocess to occur a t very acidic pH. However, the occurrence of' a similar sulfur-producing step a t neutral p H and in the presence of 0 2 has never been firmly established. T o evaluate the importance of FeS as a sorbent for SO2 removal, we have examined the aerobic conversions of FeS under a variety of SO:! and pH conditions to determine the iron products and sulfur products. Reaction of FeS with H20, with or without an SO2 source, gives rise to a weak yellowcolored solution and a finely divided precipitate, whose apparent color may irange from yellow to brown to olive to black. Most of the original FeS starting material remains unreacted (residue); however, its surface appears brassy-very much like pyrite. Physical chemical analyses, using a variety of techniques, have been performed on both the precipitate and residue fractions obtained from FeS hydrolysis. X-ray diffraction patterns (21) of FeS starting reagent, unreacted solid residue, and precipitate fractions all show the substantial presence of pyrrhotite, regardless of the conditions. (The precipitate fraction is that portion of spent slurry which remains suspended during manual swirling, but which filters as a solid. The unreacted residue fraction is that portion of the solids which remains sedimented during swirling.) Iron oxide hydrates (as FeO(0H) and, when SO2 is present, FeO) are the major constituents of the precipitate fraction resolved by X-ray diffraction under present conditions. In none of the experiments did pyrite (FeSz), elemental sulfur (SO),or any sulfur-rich iron sulfide account for more than 1%of the mass in any preparation. Bulk field-of-view X-ray fluorescence (XRF) spectra (21) of the FeS starting material, the unreacted residue fraction, and the fine precipitate fraction show decreasing S:Fe atom ratios respectively, regardless of the experimental conditions. Specifically, S:Fe atom ratios in the residue fraction are
Volume 14, Number 7, July 1980
867
2500 ppm SOz INLET GAS
I
210
1w
290
310
120
WAVELENGTH (nm)
TIME (min) 0
2
1
1
1
5
100
310
6
I
280
280
120
WAVELENGTH (nm) I 7 I4
73
70
66
63
60
56
1
I
39
29
,
,
28
21
I 2 611
SCRUBBERpH
Figure 2. Time course for titration of an aqueous scrubber with SO2 gas. Air containing 2500 ppm of SO2 was bubbled through 50 mL of solution, containing 0.05 M NaHC03and 0.2 M NH4CI,at a rate of 5 L/min. Origin on the time scale representsthe time at which SO2 was first introduced into the gas stream. (Upper Spectrum) Second-derivative ultraviolet absorption spectrum of 2500 ppm of SOz gas in air (24):scan rate, 10 nm/min; time constant, 3 s; vertical full scale, 14.9 V; modulation frequency, 45 Hz; optical path, 10 mm. Gas flow sufficiently turbulent to ensure mixing through cell. Spectra were calibrated against a sealed quartz cell containing 2.904 ppm-mol SO2 (24). (Lower Spectra) Real-time gas-phase SO2 second-derivative absorbance and solution phase pH as a function of time. Conditions are as in upper spectrum
~~~
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F Figure 1. Infrared absorption spectra of solids from hydrolysis of pyrrhotite. All spectra are recorded against a 350-mg Csl blank pellet. Procedures are as in ref 37 and 32. (A) Filtered precipitate from incubation in HzO with 0.1 M KCI; initial pH, 5.05; final pH, 5.80; pellet, 795-pg sample in 500.670 mg of Csl. (B)Filtered precipitate from incubation in 100 mL of HzOwith 30 p L of 10 M HCI; initial pH, 2.5; final pH, 5.29; pellet, 820 p g in 500.755 mg of Csl. (C) Filtered precipitate from incubation with 1.07 g of NH&I and 0.56 g of KOH in 100 mL of H20; initial pH, 9.205; final pH, 8.403; pellet, 800 p g in 499.870 mg of Csl. (D) Filtered precipitate from incubation in 100 mL of H20 with 1.72 g of Na2SZO5and 0.25 g of Na2S03;initial pH, 5.865; final pH, 3.275; pellet, 820 p g in 500.590 mg of Csl. (E) Pyrrhotite residue (unreacted) from incubation as in B above; pellet, 870 p g in 500.180 mg of Csl. (F) Lepidocrocite obtained from water tower rust, from J. J. Kovach (60); 600 p g in 498.490 mg of Csl. Mole fractions calculated for FeO(OH), FeO, or Fe(OH)*, and Fe2(S04)3-(H20)9 in all samples were based on absorbances of mineral phases in reference spectra (37,32)
868
Environmental Science & Technology
?
sulfur observed under a variety of FeS hydrolysis conditions. In the presence of NH3 a t pH 9 (C in Table I), the iron released from FeS is accounted for by iron oxide hydrates, and the sulfur by SO42-. The same is true for the hydrolysis products in dilute acid (B in Table I), except that a portion of the sulfur released from FeS is evolved as H2S gas. Substantial quantities of unidentified sulfur are observed for the hydrolysis products of FeS in H20 (A in Table I) or in bisulfite solution (D in Table I). In both cases, one can reasonably expect substantial concentrations of S(-11) and derivatives to be present min the solution phase as a result of FeS dissolution, although their form might not be known. Separate experiments in which SO2 gas is captured by solutions containing FeS give total S042- yields roughly equal to the total leached iron, as in D of Table I. However, control experiments in which FeS was omitted demonstrate that the S042-originates from the SO2 source, not from FeS (Figure 3; also ref 21). Most of the sulfur(1V) in D remains unreacted. In an attempt to identify additional sulfur-containing species as products of FeS hydrolysis in water or in bisulfites, thiosulfate ion ( S Z O ~ ~was - ) also monitored. Thiosulfate is a pivotal intermediate in the Wackenroder reaction, by which sulfides such as H2Sand FeS react with SO2 to form elemental sulfur ( 7 ) .Hence, the accumulation of S ~ 0 3 is ~ -necessary for the direct product of SO.During SO2 capture in a scrubber containing finely ground FeS, S ~ 0 3 concentrations ~are below the UV spectrophotometric detection limit (25). Conse-
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Table 1. Material Balances of End Products of Pyrrhotite Hydrolysis a Fel-xS
A
starting material mol initial mol recovered gases detected by odor solution phase pmol of Hf liberated pmol of SO4+ pmol of Fe(ll) pmol of total iron leached precipitate, mg phases identified by XRD data powder patterns bulk molar ratio !S:Fe S:Fe ratio for S-rich particles, by SEM-XRF pmol of Fe20s oir Fes04 pmol of FeO(0H:ias lep. pmol of Fe(OH)2or FeO pmol of Fe2(S04)3-(H20)9 as coq. pmol of Fel-,S pmol of FeS2 as pyrite pmol of Ss (elemental) total balances for hydrolysis nonsulfide iron, pmol sulfur, less Fel-,,S, pmol sulfur, incl. gases, pmol unidentified sulfur, pmol
B
D
C
0.100 0.102
0.100 0.100
0.100 0.104
0.100 0.099
none unbuffered
H2S
"3
so2
-1 259 10 14 79.85
-310 387 128 155 41.85
+6530* 235 0 2 21.55
+2240* 2020" 1000* 31.80
Fel-,S FeO(0H)
Fel-,S FeO(0H)
Fel-,S FeO(OH)2
Fel-,S FeO(0H) FeO
0.093, 0.133 0.8-1.1
0.107 0.9-1.1
0.193 0.9-1.1
0.29 0.9-1.1
observed
?
857
? 30 200
