Environ. Sei. Technol. 7984, 18, 726-729
ratios exceeding 6 in samples taken beneath a deciduous forest were observed. This is most likely due to the oxidation of isoprene, which is known to lead to high levels of formaldehyde and which does not produce acetaldehyde as a product (26). Acknowledgments
We acknowledge many helpful discussions with J. Gaffney. Registry No. HCHO, 50-00-0; CH,CHO, 75-07-0. Literature Cited Kitchens, J. F.; Casner, R. E.; Edwards, G. S.; Harward, W. E., 111; Macri, B. J. U.S. Department of Commerce, National Technical Information Service, 1976, Report PB-256839. Levaggi, D. A.; Feldstein, M. J. Air Pollut. Control Assoc. 1969, 19, 313. Beller, T. A.; Sigsby, J. E., Jr. Environ. Sci. Technol. 1970, 4, 150. Logan, J. A.; Prather, M. J.; Wofsy, S. C.; McElroy, M. B. JGR, J. Geophys. Res. 1981,86, 7210, Gaffney, J. S. April 1982, Brookhaven National Laboratory Report BNL-51605, UC-11, pp 1-76. Altshuller, A. P.; Miller, D. L.; Sleva, S. F. Anal. Chem. 1961, 33, 622. Lyles, G. R.; Dowling, F. B.; Blanchard, V. T. J. Air Pollut. Control Assoc. 1965, 15, 106. Altshuller, A. P.; Lent, L. J. Anal. Chem. 1963, 35,1541. Hauser, T. R.; Cummins, R. L. Anal. Chem. 1964,36,679.
Kallio, H.; Linko, R. R.; Kaitarantz, J. J. Chromatogr. 1972, 65, 355. Papa, L. J.;Turner, L. P. J. Chromatogr. Sci. 1972,10,744. Hoshika, Y.; Takata, Y. J. Chromatogr. 1976, 120, 379. Selim, S. J.Chromatogr. 1977, 136, 271. Kuwata, K.; Vebori, M.; Yamasaki, Y. J. Chromatogr. Sci. 1979, 17, 264. Kuntz, R.; Lonneman, W.; Namie, G.; Hull, L. A. Anal. Lett. 1980, 13, 1409. Fung, K.; Grojean, D. Anal. Chem. 1981,53, 168. Lowe, D. C.; Schmidt, U.; Ehhalt, D. H. Geophys.Res. Lett. 1980, 7, 825. Lowe, D. C.; Schmidt, U.; Ehhalt, D. H.; Frischkorn, C. G. B.; Nurnberg, H. W. Environ, Sci. Technol. 1981,15,819. Grosjean, D.; Fung, K. Anal. Chem. 1982,54, 1221. Beasley, R. K.; Hoffman, C. E.; Rueppel, M. L.; Worley, J. W. Anal. Chem. 1980,52, 1110. Neitzert, V.; Seiler, W. Geophys. Res. Lett. 1981, 8, 79. Grosjean, D. Environ. Sci. Technol. 1982, 16, 254. Carter, W. P. L.; Winer, A. M.; Pitts, J. N., Jr. Atmos. Environ. 1982, 16, 113. Calvert, J. G.; Stockwell, W. R. Environ. Sci. Technol. 1983, 17, 4286. Gaffney, J.;Tanner, R. L., Brookhaven National Laboratory, Upton, NY, unpublished data, 1983. Cox, R. A.; Derwent, R. G.; Williams, M. R. Environ. Sci. Technol. 1980, 14, 57.
Received for review November 14,1983. Accepted March 26,1984. This work was generously supported by the Officeof Health and Environmental Research, US.Department of Energy, and performed under Contract DE-AC02-76CH00016.
Iodomethane as a Potential Metal Mobilizing Agent in Nature? John S. Thayer * Department of Chemistry, University of Cincinnati, Cincinnati, Ohio 4522 1
Gregory J. Olson and Frederlck E. Brlnckman Inorganic Materials Division, National Bureau of Standards, washington, DC 20234
Iodomethane, an ubiquitous biogenic metabolite, has been found to release metals into water from both polluted anoxic sediments and certain metal compounds. Reactions of water-insoluble binary and ternary sulfides with iodomethane gave methylsulfur compounds as determined by flame photometric gas chromatography and proton nuclear magnetic resonance spectrometry. Naturally occurring iodomethane may react with metal sulfides or metals under certain environmental conditions to generate water-soluble and/or volatile derivatives. Introduction
Biological processes, both exo- and endocellular, affect the environmental abundance and redistribution of many elements (I). Such mobilizing processes have commercial value in metal ore solubilization (2) but may also cause adverse environmental impacts through toxic metal mobilization and methylation ( 3 , 4 ) . One such extracellular metabolite is iodomethane, CH31,found throughout the earth's atmosphere (5), especially over marine waters, 'Portions of this work were presented at the 186th National Meeting of the American Chemical Society, Washington, DC, Sept 1983; ENVR 170. t National Science Foundation Faculty Research Participation Fellow, 1983; National Bureau of Standards Guest Worker, 1983. 726
Environ. Sci. Technol., Vol. 18, No. 9, 1984
where it is generated by seaweeds (6). This compound has been proposed as a contributor to the corrosion of iron (7) or copper (8) by seawater and as a reactor with metallic lead to form tetramethyllead (9). During an investigation into the formation of methylmetal compounds in sediments, we found that iodomethane added to samples of anoxic sediments from rivers and bays generated dimethyl sulfide and markedly increased the dissolution of several metals. Tin, lead, and mercury are among those elements previously believed to be immobilized in natural sediments. Increasing commercial production of organotin compounds has resulted in these materials entering the environment at global levels comparable to tin levels introduced by natural weathering (IO). While most of these metal compounds probably end up in estuarine or marine sediments (Il), their chemical forms and persistence are presently unknown (12). We have investigated the effect of iodomethane on compounds of various metals and metalloids to see if CH31might serve as a mobilizing agent and perhaps be a useful reagent for the study of metal compounds bound in sediments. Experimental Section
Chemical and Chemical Procedures. Lead chalconides, metal sulfides, minerals, and iodomethane were obtained from commercial suppliers. Enargite (Cu3AsS4)and
0013-936X/84/09 18-0726$01.50/0
0 1984 American Chemical Society
arsenopyrite (FeAsS) were obtained as mineralogical specimens and were ground to sub-40 mesh particles. Lead sulfide, mercuric sulfide (black), and stannous sulfide were prepared by addition of sodium sulfide to solutions of Pb(N03)2,HgC12,and SnC12,respectively. The first two sulfides were washed several times with deionized water and dried at 60 "C. Stannous sulfide was rinsed with anoxic deionized water and stored under N2 without drying. In volatile methylsulfur evolution studies, and in certain metal sulfide solubilization studies, 100 mg of metal sulfide and 3 mL of deionized water were placed into 5-mL glass vials. After 15-min contact, the water was siphoned off (to remove any soluble metal species), and 1.0 mL of deionized water and 5 pL of CHJ were added. The vials were sealed with Teflon septum screw caps and were then incubated at 22 "C in the dark. After 3 days headspace samples for gas chromatographic analysis were removed by syringe. For most solution studies involving metal sulfides or sediments, 500-800 mg of solid was placed into a Whatman cellulose extraction thimble, which was subsequently put into 15 mL of deionized water in a screwtop vial. Iodomethane (25-50 pL) was added directly into the thimble and the vial sealed with a screwtop cap. Subsequent aliquots were removed from the water outside the thimble for atomic absorption spectrophotometry. Gas Chromatographic Procedures. Volatile sulfur compounds were detected by injecting vapors from experimental mixtures into a gas chromatograph equipped with a flame photometric detector (GC-FPD) (13). Vapor components were separated on a 6-ftChromosil330 Teflon column (Supelco, Inc., Bellefonte, PA) with an inner diameter of 1/8-in. The column temperature was maintained at 25 "C for the first 4 min and then heated at 32 "C min-l to 60 "C. Zero-grade N2 was used as carrier gas, at a flow rate of 20 mL min-l. The flame photometric detector operated with flame gas flows of H2 (50 mL min-l), air (50 mL min-l), and O2 (10 mL m i d ) and with a detector temperature of 200 "C and injection port temperature of 150 "C. A selective interference filter (394 nm) was used for sulfur studies. Methyltin compounds were detected by using tin-selective purge and trap GC-FPD (14). Aqueous samples were treated with sodium borohydride to generate volatile methylstannanes (14). Most peaks in analytical samples were identified by comparison of retention times to those for authentic standards. Atomic Absorption Procedures. Commercially available atomic absorption spectrophotometers, employing hollow cathode or electrodeless discharge lamps, were used for quantitation of metals in reaction solutions. Calibration curves were prepared for each run with standard metal solutions. Nuclear Magnetic Resonance Procedures. Fourier transform nuclear magnetic resonance (FT-NMR) spectrometers operating at 200 or 400 MHz for lH were used to study heterogeneous mixtures of diamagnetic, particulate metal sulfides and CH31 in D20.
Results When samples of anoxic sediment from Baltimore Harbor were treated with iodomethane and aliquots of the supernatant water analyzed by purge-and-trap GC-FPD (14), appreciable quantities of dimethylstannane and dimethyl sulfide appeared almost immediately; tetramethyltin and dimethyl disulfide subsequently appeared after overnight standing. Similar studies using atomic absorption spectrometry showed that after 24 h iodomethane had caused dissolution of other metals from
Table I. Release of Metals from Baltimore Harbor Sedimentsn Colgate Creek Pb Mn Cu sediments/CHJ sediments only
1.360 0.460
Pb
0.365 0.032 0.268 0.017
Jones Falls Mn Cu
1.510 0.770
0.029 0.015
0.056 0.013
Numbers represent levels of dissolved metals in milligrams per liter. Table 11. Dissolution of Lead from Lead Chalconides" contact times, ks PbS 8.24 15.75 77.70
PbS/ CHJ 0.70 4.91 17.50
0.04 0.11 0.51
PbSe/ CHJ
PbSeb 0.05 0.04 0.01
4.11 4.89 14.50
PbTeb 0.13 0.07 0.01
PbTe/ CH31 4.47 7.03 15.40
ONumbers represent levels of dissolved lead in milligrams per liter. "The apparent decrease in soluble lead probably reflects analytical uncertainties at such low concentrations. Table 111. Dissolution of Iron, Mercury, and Arsenic from Sulfides in Water
reactants
time, days
metal released
FeS + CHJ FeS only FeS, + C H J FeSi only HgS (red) + CH31 HgS only HgS (black) + CHJ HgS only enargite" + CHJ enargite only arsenopyriteb + CH31 arsenopyrite only
1
Fe
1
Fe
4
Hg
4
Hg
1
As
1
As Fe
so1ution concn, mg
L-' 155 0.1 22.0 9.8 1.9 0.05 2.7 0.04 1.36 0.57 0.63 0.38 18.0 1.2
Enargite, CusAsSa. *Arsenopyrite, FeAsS.
Baltimore Harbor sediments (Table I). Samples of sediments from less industrially polluted portions of the Chesapeake Bay showed little or no CHJ-stimulated Cu or P b release. Iodomethane added to samples of sediments from a metal-polluted portion of the Rhine River in West Germany (15) caused a 3-fold increase in copper release compared to CH31-freecontrols after 24-h standing. The relationship of sediment metal loading to CHJ-stimulated metal release from sediments requires additional study, especially with respect to the speciation of metals in sediments. A strong odor of hydrogen sulfide appeared when Baltimore Harbor sediments were treated with strong acids, suggesting that the metals occurred in the sediments predominantly as the sullides and that they were dissolving primarily through reaction with iodomethane to form methylsulfur compounds. To test this hypothesis, we investigated the reaction of metal sulfides in deionized water with iodomethane. Lead sulfide, and other lead chalconides, showed greatly enhanced water solubility in the presence of iodomethane (Table 11). Pb(SCH3), also reacted with iodomethane to form dimethyl sulfide. FeS, FeS2, HgS, enargite, arsenopyrite (Table 111),and MnS showed similarily enhanced solubilities. Comparative studies using iodomethane and various authentic metal sulfides (Table IV) indicated that methylsulfur compounds Envlron. Scl. Technol., Vol. 18, No. 9, 1984
727
PROPOSED MECHANISM FOR
Table IV. Relative Quantities of Released Methylsulfur Compounds"
THE METHYLATION OF SOLID METAL SULFIDES
metal sulfide Cas PbS FeS AsSz SnSz HgS (red)
(CHAS
(CH3)zSz
3 000 000 110000 55 000
448000
50 000
14000 7 700
1000 2 400
'S
" Numbers represent integrated chromatographic areas in arbitrary units. SnS2
.,
-
PbS
/
Figure 2. Proposed mechanlsm for the reaction of iodomethane with Insoluble metal sulfides.
-
Discussion Iodomethane has long been known to react with elements of groups 5A and 6A, causing oxidative addition: R,E:
Me2S
Me2S2
Me1 blank
M(SCH& I
I
I
4
S
MIN
Flgure 1. Volatile methylsulfur compound detection by flame photometric gas chromatography. Headspace gas (0.1-1.0 mL) above metal sulfide-CH,I mixtures in sealed gas vials was analyzed and found to contaln methylsulfur specles. Gas above an iodomethane blank gave no detectable signal in the sulfur-selectlve analytical mode. No methylsulfur species were detected above metal sulfides without iodomethane.
were formed at varying rates of reaction. No methylsulfur compounds could be detected above metal sulfides not treated with iodomethane. Gas chromatograms of representative experiments appear in Figure 1. A mixture of gaseous H2Sand gaseous CH31showed no reaction after 20 h. By contrast, a mixture of dilute CH31 and NazS in DzO showed very rapid reaction at 27 "C; proton NMR spectroscopy indicated that reaction was ). complete in minutes with a k2 of 0.232 L mor1 ~ ~ ( 1 6The solution became more acidic as the reaction progressed. Methyl mercaptan also reacted with iodomethane in water to give dimethyl sulfide. Iodomethane has been reported to react with metals in the presence of water (7-9). Three groups have reported that P b metal formed (CH3)4Pbin the presence of CH31 (9, 17, 18). Metallic tin and stannous salts produce (CH3),Sn (17). We found that 162 mg of 30-mesh Fe powder 728
Environ. Sci. Technol., Voi. 18, No. 9, 1984
+ CH31
-
R,ECH3+
+ I-
(1)
Many metals having an available pair of electrons can react with CH31 to form methylmetal compounds under anhydrous conditions (19). Iodomethane reacts rapidly with aqueous sulfide ion to form dimethyl sulfide, which reacts more slowly to form trimethylsulfonium ion (16,20). We propose the mechanism shown in Figure 2 for the reaction of water-insoluble metal sulfides with CH31. The sulfur atoms in such systems have a large degree of covalent character and serve as brides between metal atoms to form large polymeric arrays. Oxidative addition by iodomethane would break such bridges, at sulfur or metal centers, weakening the lattice and eventually allowing the metal to dissolve. If intermediate metal thiomethoxides form, these would also react with iodomethane. Where a stable methylmetal species is not expected, as with transition metals and lead, the reaction may be written as
Me1
0
/
in contact with distilled water for 4 days gave less than 0.1 mg 1-l dissolved iron, but with 200 NLof CH31added, the level of dissolved iron increased to 166.7 mg L-l over the same time period. Both P b and Cu metals also showed increased water solubility under the same conditions.
~ n ~ 2 + M e1
1
'S
'S
/
-
+ 2CH31
M2++ 21-
+ 2(CH3)2S
(2)
We detected no methyllead species by NMR but did observe (CH3&3 and (CH3)3S+ion in an aqueous CH31PbS mixture. We also found that Pb(SCH3)2,a likely intermediate in this mechanism, reacted with CH31 in water to produce (CH3)2S. With tin, a successful oxidative methylation occurs to form a Sn(1V)-C bond (16),written approximately SnS + 3CH31 CH3Sn13+ (CH3),S (3)
-
As in the lead system, we suspect that SII(SCH~)~ may form as an intermediate, but analogous simple intermediates may not prevail with other metals. In either reaction 2 or 3, nucleophilic attack on sulfur centers in polymer bridges begins first. Our results are consistent with the possibility that naturally occurring iodomethane might react with metal sulfides present in anoxic sediments or finished metals in human artifacts in aquatic systems, thereby creating water-soluble and/or volatile derivatives. The strong reducing powers of anoxic environments will cause any metal or metalloid derivatives present to be in the mostly highly reduced form possible (21); this has been reported for selenium in marine environments (22). The strong reducing powers of anoxic sediments and the rather low reduction potentials for tin and lead suggest that, under
some conditions, these might occur in sediments as the free metals. If so, reaction with iodomethane would cause them to redissolve. The concentration of iodomethane in the environment will have an important effect on the extent and rates of metal solubilization. Although data are sparse, the average iodomethane concentration in seawater is about 1.0 ng L-l (23). However, in areas of high biomass productivity, iodomethane concentrations can be loo0 times higher (24). To our knowledge, the levels of iodomethane in sediments or pore waters have not been reported. However, the occurrence of iodomethane in some sediments seems likely since dimethylsulfonium compound decomposition in iodine-accumulating algae is a proposed route of iodine methylation (25). The reaction of iodomethane with dissolved or metalbound sulfide represents a potential source of environmental dimethyl sulfide. We do not view this reaction as a major contributor to overall environmental levels of metals or dimethyl sulfide since the dimethyl sulfide concentration in the ocean is 2 orders of magnitude higher than that of iodomethane (26), and 2 or more mol of iodomethane is required to produce dimethyl sulfide from metal-bound or dissolved sulfides. More important sources of dimethyl sulfide probably include the aforementioned decomposition of dimethylsulfoniumcompounds in marine algae and phytoplankton (25) and microbial degradation of sulfur-containing amino acids (27). While there is presently insufficient evidence to state that iodomethane enhances the overall circulation of metals throughout the environment, the evidence does point to that strong possibility in certain local environmental compartments. Especially noteworthy is the suggestion that biogenesis of iodomethane itself is a concurrent process in dimethyl sulfide production in algae (25). More data on iadomethane production and occurrence in sediments are needed for better assessment of its role in metal cycling. There is ample scope for further investigation into this area.
Acknowledgments We thank Thomas W. Brueggemeyer, University of Cincinnati, for running the atomic absorption spectra and Joseph A. Caruso, University of Cincinnati, for the use of his atomic absorption spectrometer. We also thank William Manders, University of Maryland, for assistance with the nuclear magnetic resonance spectroscopic investigations. Registry No. Cu&S4,14933-50-7; FeAsS, 1303-18-0; (CH3)a2, 624-92-0; CHsI, 74-88-4; Fe, 7439-89-6; Hg, 7439-97-6; As, 744038-2; Pb, 7439-92-1; Sn, 7440-31-5; Ca, 7440-70-2.
Literature Cited (1) Ehrlich, H. L. “Geomicrobiology”; Marcel Dekker: New York, 1981. (2) Brierley, C. L. CRC Crit. Rev. Microbiol. 1978, 6, 207.
(3) Thayer, J. S.; Brinckman, F. E. Adv. Organomet. Chem. 1982, 20, 313. (4) Thayer, J. S. “Organometallic Compounds and Living Organisms”;Academic Press: New York, 1984; pp 216-245. (5) Singh, H. B.; Salas,L. J.; Stiles, R. E. Environ. Sci. Technol. 1982, 16, 872-880. (6) Lobban, C. S.; Wynne, M. J.; Eds. “The Biology of Seaweeds”; University of California Press: Berkeley, 1981; p 608. (7) More, E. D.; Beccaria, A. M. h o c . Int. Congr. Met. Corros. 7th 1979, 1324-1332. (8) More, E. D.; Beccaria, A. M. Br. Corros. J. 1977,12,243-246. (9) Snyder, L. J.; Bentz, J. M. Nature (London) 1982, 296, 228-229. (10) Byrd, J. T.; Andreae, M. 0. Science (Washington, D.C.) 1983,218, 565-567. (11) Turekian, K. K. Geochim. Cosmochim. Acta 1977,41,1139. (12) Brinckman, F. E.; Olson, G. J.; Ivenon, W. P. “Atmospheric Chemistry”;Goldberg, E. D., Ed.; Springer-Verlag: Berlin, 1982; pp 231-249. (13) Olson, G. J.; Brinckman, F. E.; Jackson, J. A. Int. J . Environ. Anal. Chem. 1983, 15, 249. (14) Jackson, J. A.; Blair, W. R.; Brinckman, F. E.; Iverson, W. P. Environ. Sci. Technol. 1982, 16, 110-119. (15) Iverson, W. P.; Tobschall, H. J.; Brinckman, F. E.; Olson, G. J., paper presented at the 6th International Symposium on Environmental Biogeochemistry, Santa Fe, NM, Oct 10, 1983. (16) Manders, W. F.; Olson, G. J.; Brinckman, F. E.; Bellama, J. M. J . Chem. SOC.,Chem. Commun. 1984,538. (17) Craig, P. J.; Rapsomanikis, S. J . Chem. SOC.,Chem. Commun. 1982,1982, 114. (18) Jarvie, A. W.; Whitmore, A. P. Environ. Technol. Lett. 1981 2, 197. (19) Rochow, E. G. J . Chem. Educ. 1966,43, 58. (20) Reid, E. E. “Organic Chemistry of Bivalent Sulfur”; Chemical Publishing Co., Inc.: New York, 1960; Vol. 11. (21) Kochi, J. K. In “Organometals and Organometalloids. Occurrence and Fate in the Environment”; Brinckman, F. E.; Bellama, J. M., Eds.; American Chemical Society: Washington, DC, 1978; pp 205-234. (22) Cutter, G. A. Science (Washington, D.C.) 1982, 217, 829-831. (23) Singh, H. B.; Salas, L. J.; Stiles, R. E. J . Geophys. Res. 1983, 88, 3684-3690. (24) Lovelock, J. E. Nature (London) 1975,256, 193-194. (25) White, R. H. J . Mar. Res. 1982, 40, 529-536. (26) hdreae, M. 0.;Raemdonck, H. Science (Washington,D.C.) 1983, 221, 744-747. (27) Bremner, J. M.; Steele, C. G. Adv. Microb. Ecol. 1978, 2, 155-201.
Received for review January 3,1984. Accepted March 19,1984. This research was supported in part by the Office of Naval Research. J.S.T . thanks the National Science Foundation for an IRP Grant. Certain commercial equipment and materials are identified in this paper in order to adequately specify the experimental procedures. In no case does such endorsement imply recommendation or endorsement by the National Bureau of Standards, nor does it imply that the material or equipment identified is necessarily the best available for the purpose.
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