The simultaneously extracted

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1 Chemical Dynamics of SEDIMENTARY Acid Volatile Sulfide USFWS JOHN W. MORSE TEXAS A&M UNIVERSITY DAVID RICKARD CARDIFF UNIVERSITY (WALES) JOHN MORSE These measurements have implications for metal toxicity but lack consistent definition. The simultaneously extracted metal acid volatile sulfide (SEM AVS) method is widely used to estimate whether or not metals in sediments are toxic. Although considerable attention has been given to metal sulfide interactions and related toxicological research, the complexities of the sedimentary sulfide system and how they may influence the SEM AVS relationship have received scant consideration. 24 American Chemical Society APRIL 1, 24 / ENVIRONMENTAL SCIENCE & TECHNOLOGY 131A

2 In reviewing the many published AVS analyses of sediments, we were struck by the extreme variability in the results. The method rests on the basic concept that many toxic metals, such as cadmium, lead, zinc, mercury, and copper, readily precipitate as highly insoluble metal sulfide minerals when exposed to dissolved sulfide, S( II). But what do the AVS and SEM values obtained from sediment analyses actually represent? Both are operational definitions. AVS and SEM values are obtained by adding acid to the sediment and measuring the amount of S( II) evolved and the metal concentrations in the solution. We refer to the products of this analysis as AVS for brevity. The operationally defined volatile sulfides comprise a group of what are generally believed to be metastable iron sulfide minerals and dissolved S( II) species that, when exposed to HCl, form H 2 S which can subsequently be collected and analyzed. These volatile sulfides are presumed to provide a metalreactive source of S( II). Thus, when the AVS concentration in sediments exceeds that of the metals under consideration, the metals precipitate and should not be toxic. Initial studies with organisms supported this concept (1, 2). However, the concept is confusing. Uncomplexed aqueous S( II) is highly toxic to most organisms and even sulfate-reducing bacteria, which produce almost all sedimentary S( II) and have a limited free S( II) tolerance. More recent toxicological studies have raised serious questions about the general applicability of the earlier work (3, 4). In other words, toxicity is not only a function of the metal concentrations; the toxic effects of metals and S( II) need to be deconvoluted as well. With occasional exceptions, AVS is confined to a relatively small portion (<1%) of total inorganic sedimentary sulfides (total reduced sulfide [TRS]). AVS is usually ephemeral, often forming in the top few centimeters of anoxic sediments but decreasing with depth to undetectable concentrations within 2 centimeters (cm) or less. Although AVS commonly represents only a small portion of TRS, it has received considerable attention because these volatile sulfides are one of the most chemically reactive components of anoxic sediments during early diagenesis. The primary purpose of this paper is to discuss the rather diverse literature on sedimentary AVS. We approach the problem by considering the chemical dynamics of AVS formation and loss in sediments because of its direct bearing on the role AVS plays in sequestering toxic metals and dissolved S( II) from organisms. Definition and chemistry of AVS The problem with an operationally defined concept such as AVS is that it is procedure-dependent, and the reported procedures have varied among investigators (5). Berner defined AVS as the sulfide that evolved from sediments after enough 1. N HCl was added to cover a weighed sediment sample in a flask and the sample was heated to boiling, thereby expelling all of the sulfide (6). This approach measures both solid-phase HCl-reactive sulfide minerals and S( II) dissolved in pore waters. Some researchers have assumed that AVS is associated entirely with iron sulfide minerals (7 ). However, this is a dangerous assumption because AVS may be almost all dissolved S( II). It also raises the question of whether solid- and dissolved-phase AVS fractions have potentially very different influences on trace metal behavior. More recent studies show that the dissolved component includes aqueous FeS clusters (8). These clusters are defined as groups of molecules that are intermediate to the bulk solid and aqueous ions and that form dissolved, sometimes aqueous, groups of molecules and nanoparticles. Experimental studies using synthetic iron sulfide minerals indicate that although HCl may completely dissolve FeS, it only partially dissolves Fe 3 (greigite), and it can dissolve a variable small fraction of FeS 2 (pyrite) that is similar to the fraction of TRS often reported as AVS (9). Experimental conditions and extraction time influence results. In sediments, the sulfide minerals extracted by the AVS procedure may vary between very early diagenesis, when iron sulfide mineral formation is most active, and more mature sediments, when grain size and degree of crystallinity increase. AVS in marine sediments and waters With rare exceptions, AVS minerals are not directly observable by traditional techniques such as X-ray diffraction analysis and scanning electron microscopy (1). Our experience is that they are often difficult to detect even with high-resolution electron microscopy. Their presence is inferred by the previously discussed chemical leaching techniques and also by comparing ion activity products to the apparent equilibrium solubility products of the minerals. However, although AVS has been commonly measured, AVS minerals have not generally been the focus of sulfidic sediment studies. Robert Aller s dissertation remains the most comprehensive study of AVS behavior in marine sediments and has become the basis for many of the widely held ideas about formation and loss in anoxic marine sediments (11). The classic profile is shown for the FOAM site in Long Island Sound in Figure 1a. TRS increases rapidly down to about 4 cm and then is close to constant. AVS, a small fraction of TRS, has a maximum value at about 4 6 cm and decreases to almost undetectable values at greater depths. A relationship not often noted is that the AVS maximum is at the base of the zone of TRS formation. It is also close to coincident with the depth at which dissolved iron decreases to low values and detectable dissolved S( II) rapidly increases. The intrinsic weakness of any reported AVS val- 132A ENVIRONMENTAL SCIENCE & TECHNOLOGY / APRIL 1, 24

3 FIGURE 1 Acidic volatile sulfide (AVS) and total reduced sulfide (TRS) depth profiles Data vary greatly by location. AVS data are in red, and TRS data are in blue for (a) the FOAM site in Long Island Sound (11), (b) the Louisiana shelf west of the Mississippi River delta (12 ), and (c) Laguna Madre, Texas (13 ). Units are micromoles per gram dry weight, and the black arrow indicates the zone of major sulfide formation. (a) Depth (cm) (b) Depth (cm) (c) Depth (cm) AVS and TRS ( mol/gdw) AVS and TRS ( mol/gdw) AVS and TRS ( mol/gdw) ues from an individual site is the geographic limitation, which calls into question the generality of the findings. In reviewing the many published AVS analyses of sediments, we were struck by the extreme variability in the results. For example, studies of sulfide geochemistry throughout much of the Gulf of Mexico and some of its adjacent estuaries demonstrated a wide variety of patterns of AVS distributions (12, 13) and few resembled those in Long Island Sound. In many cores, no detectable AVS was observed, but in others it was the major TRS component, which increased with depth (Figures 1b and c). It has also been observed that AVS comprises a major and persistent fraction of TRS in both iron-rich (14) and iron-poor sediments (13), shallow water sediments, and sediments from euxinic (anoxic and sulfidic) marine environments (15). In contrast, other modern sediments containing substantial amounts of TRS with moderate to high sulfate reduction rates have nearly undetectable concentrations of AVS, and pore waters apparently are undersaturated with respect to mackinawite (tetragonal FeS; 16). These observations indicate that there is no correct depth at which to sample for AVS. The AVS approach can yield quite different results at various depths, even relatively near the sediment water interface. Aller also observed considerable seasonable variability of AVS in the top 3 cm of sediment. Major decreases occurred in the winter (11), and Aller hypothesized that because of the colder temperatures, FeS formation rates were lower and a smaller supply of FeS-rich particles were brought up from below by bioturbation. Recent Chesapeake Bay sediment studies provide another example of greater than order of magnitude variation in AVS seasonal concentrations (17). A separate study in the Chesapeake Bay indicated that the seasonal variability in AVS, when AVS was <2% of HCl-extractable iron, could influence the flux of dissolved copper across the sediment water interface and the concentrations of copper and zinc in nearinterfacial sediments (18). These results clearly illustrate that the AVS pool in sediments is very dynamic and that the extent to which AVS may limit the toxicity of metals is probably seasonally dependent. We observed that it is difficult to make a priori predictions of AVS concentrations, depth distributions, and seasonal patterns from major sediment attributes. This also holds true for the little-studied relationship between the dissolved and solid-phase subcomponents of AVS. If AVS does indeed exert a major influence on metal toxicity, then these far from trivial complexities in sulfide distribution and temporal variability are certainly of first order importance to understanding the impact of toxic metals on benthic ecosystems. AVS components The solid AVS phases are believed to consist predominantly of iron sulfides, because iron is by far the most abundant transition metal in the majority of these environments. The low-temperature iron sulfide minerals include mackinawite, greigite (cubic Fe 3 ), troilite (hexagonal FeS), the pyrrhotites (hexagonal APRIL 1, 24 / ENVIRONMENTAL SCIENCE & TECHNOLOGY 133A

4 Clearly, AVS is far from a simple component of anoxic sediments whose occurrence and behavior we can readily predict at this time. FIGURE 2 Fe 1 x S), pyrite (cubic FeS 2 ), and marcasite (orthorhombic FeS 2 ). All of these phases have been reported from sedimentary environments (18). In addition, smythite (hexagonal Fe 3 ) and cubic FeS occur in low-temperature aqueous systems but have not, as yet, been detected in sedimentary systems. By far, the most abundant solid components of AVS are believed to be metastable mackinawite and greigite. However, as previously noted, a minor fraction of pyrite may in some instances comprise the dominant AVS solid phase. Structure of mackinawite Mackinawite is characterized by sheets of Fe Fe atoms with an Fe Fe bond suggesting extensive Fe Fe bonding. Nanoparticulate FeS contains H 2 O in layers between the sheets and in the interstices between the sulfur atoms, and possibly many other environmental components could be included in these sites. Nanoparticulate mackinawite displays a platy habit, but more than 5% of the surface is characterized by edge positions as against typical {1} surfacedominant bulk material. This means that the sorption characteristics of the nanoparticulate material are likely to be considerably different from the bulk mineral..5 nm Mackinawite is believed to be a substantial component of AVS because it has been detected in sedimentary environments (18) and it is kinetically the first form of iron sulfide to precipitate from most lowtemperature aqueous systems (21). Greigite may be a minor component of AVS because it has been detected in sedimentary environments (18). It is highly ferrimagnetic and has been known to contribute to the remnant magnetism of sediments, notably in lakes. However, all this evidence is indirect. There is no case in which AVS concentrations have been directly related to mackinawite and greigite concentrations. Also, as previously noted, probably only a nonquantifiable portion of greigite and pyrite is extracted by the AVS procedure. Because the pk 1 for H 2 S is close to 7, the major dissolved S( II) components of AVS are generally assumed to be H 2 S (aq) and HS, depending on the system s ph. The pk 2 for H 2 S is extremely high, which suggests that the aqueous S 2 ion is not a major contributor to AVS in the natural environment. In addition, many other more oxidized dissolved sulfur species, including polysulfides and the intermediate sulfuroxy anions, are normally not significant species in the total (S II) budget. Furthermore, dissolved iron complexes may contribute to AVS. FeS clusters have been implicated in sedimentary pyrite formation (8). These clusters are suggested to have a basic one-to-one stoichiometry (22) and range from simple, dissolved molecular clusters to nanoparticles. On the basis of these considerations, we offer the following equations for the concentration of AVS in sediments: Fe S C AVS = C S( II)dissolved + C S( II)minerals + C S( II)clusters and nanoparticles (1) C S( II)dissolved = C H2 S (aq) + C HS + C S x + C FeHSx (2) AVS is often referred to simply as FeS. Recent work has established that FeS is mackinawite and that phases such as amorphous FeS do not exist (18). The first condensed phase in the system is hydrated nanoparticulate mackinawite with a particle size of 2 nanometers, and the precursor aqueous FeS clusters have mackinawite-like structural elements (2). Over time, nanocrystalline mackinawite dehydrates, grows in size, and acquires the properties of the bulk phase (Figure 2). However, the nanocrystalline component has varying solubility. During AVS measurements, nanoparticulate mackinawite will not be retained as a solid phase by conventional pore water filtration with, for example,.45-micrometer filters. This means that indirect attempts to prove that mackinawite constitutes a substantial component of AVS using solubility considerations suffer considerably from both theoretical and practical uncertainties. C H2 S, mineral = C FeS + C ƒfes2 (3) where S x represents oxidized dissolved sulfur species that evolve H 2 S following reaction with HCl; FeHS x refers to Fe(II) sulfide complexes; FeS is the sum of the phases, like mackinawite and greigite, that react with HCl to produce H 2 S; and ƒfes 2 is the fraction of pyrite that dissolves in HCl. Chemical dynamics of AVS Clearly, AVS is far from a simple component of anoxic sediments whose occurrence and behavior we can readily predict at this time. The heart of the problem is that AVS is a complex mixture of metastable dissolved and solid phases. AVS gives only a partial representation resulting from the incomplete analysis of the reaction intermediaries typically arising during the conversion of microbiologically produced sulfide 134A ENVIRONMENTAL SCIENCE & TECHNOLOGY / APRIL 1, 24

5 to stable pyrite. Knowing the factors controlling the kinetics of the interactions is thus necessary for understanding sedimentary AVS behavior (Figure 3). Mackinawite is a metastable phase with respect to pyrite and pyrrhotite in sedimentary environments (Figure 4). However, it precipitates rapidly (<1 millisecond) from aqueous solutions at ambient temperatures, whereas pyrite nucleation is relatively slow. The relative instability of mackinawite with respect to pyrite has led to some confusion in the literature regarding the relationship between the two phases. During equilibration, pyrite will typically replace mackinawite as the dominant phase in the system. However, this replacement is not a literal process because equilibrium is related to the thermodynamic properties of a defined system. And the boundaries of the system being considered are usually far larger than an individual mackinawite nanoparticle. This equilibration process merely means that within the bounds of the system, pyrite will form and mackinawite will dissolve over an unspecified time. It does not imply that mackinawite is physically converted to or replaced by pyrite in situ. Mackinawite forms via two pathways involving reactions with H 2 S and HS (21). The formation mechanism is therefore ph-dependent. Greigite, the thiospinel of iron (Fe 3 ), has not been reported to form directly from solution. All syntheses of this material require preexisting mackinawite. Detailed electron diffraction studies of the mackinawite-to-greigite transition demonstrate that this is a solid-state transformation, with a structural congruency between the two phases (23). The reaction is highly sensitive to catalysts and inhibitors. Pyrite can be synthesized by a virtually infinite number of recipes that may have little or no relationship to the pyrite-forming processes in sedimentary environments. Likewise, trying to extract mechanistic information from sedimentary analyses is extremely difficult. The fundamental chemical mechanisms for pyrite formation can only be determined through rigorous experimental kinetic studies. And these mechanisms lie at the heart of all the various synthetic recipes and sedimentary pyrite-forming processes. Pyrite is known to form via two mechanisms. In suboxic environments, the rate of pyrite formation is a function of the concentration of polysulfides, S x ( II), in solution (24). Because pyrite is an iron(ii) disulfide, the result is unsurprising. However, inorganic polysulfides are generally not abundant in anoxic sedimentary environments, and their availability for this reaction is likely to be rate-limiting. The kinetics of the reaction between mackinawite and aqueous H 2 S in anoxic environments were shown to have significant rates even at low temperatures (25). The mechanism involves the oxidation of FeS by H 2 S, but not HS. Interaction of trace metals with sedimentary sulfides Many metals and dissolved S( II) are rendered nontoxic to organisms by AVS because of the low solubility of their metal sulfide minerals. Thus, toxicity FIGURE 3 Factors controlling pyrite s kinetics of interactions Simplified kinetic schemes can summarize present knowledge of the chemical dynamics of sedimentary pyrite formation in high- and lowsulfidation systems in suboxic and anoxic environments. The iron sulfide total ion activity product is IAP, and K s FeS m represents the local solubility coefficient for mackinawite. Red equations indicate reactions in suboxic environments, whereas blue marks anoxic conditions. (a) In suboxic environments, aqueous S( II) species are oxidized by electron acceptors such as S(), Fe(III), and O 2, with at least partial microbiological mediation, to produce polysulfides, S x ( II). (b) In high-sulfidation environments, dissolved Fe(II) species react with aqueous S( II) species to form aqueous FeS clusters, FeS aq, which condense to form mackinawite. In suboxic environments, the clusters react with polysulfides S x ( II), but in anoxic environments they combine directly with H 2 S to form pyrite. (c) In lowsulfidation environments, mackinawite dissolves to produce aqueous Fe(II) and S( II) species. Depending on conditions, dissolved Fe(II) species can then react with polysulfides or, apparently, with H 2 S to produce pyrite. (d, e) Different pyrite textures commonly produced in these different systems. The pyrite texture, in terms of pyrite crystal size, is related to the relative rates of nucleation and growth, which in turn is related to the various degrees of supersaturation with respect to pyrite usually encountered in these systems. (a) S( II) + e S x ( II) (b) High-sulfidation environments (IAP > K s FeS m ) Fe(II) + S( II) FeS aq FeS m FeS aq + S x ( II) FeS 2 FeS aq + H 2 S FeS 2 (c) Low-sulfidation environments (IAP < K s FeS m ) FeS m Fe(II) + S( II) Fe(II) + S x ( II) FeS 2 Fe(II) + H 2 S FeS 2 (d) High-sulfidation environments (IAP > K s FeS) Nucleation > Crystal growth Nucleation > Crystal growth Microcrystalline pyrite Microcrystalline pyrite (e) Low-sulfidation environments (IAP < K s FeS) Nucleation > Crystal growth Nucleation < Crystal growth Microcrystalline pyrite Pyrite single crystals and overgrowths is reduced when the AVS concentration is greater than the concentration of the simultaneously extracted metals. However, because these metals generally occur at relatively trace concentrations, even in contaminated sediments, nucleation of the pure phases is probably unlikely. It is more likely that these trace concentrations are removed via coprecipitation or metal-exchange reactions with the more dominant iron sulfide phases (26). The sorption properties of nanocrystalline mackinawite are very different from those of the bulk, and they change with time. As a consequence, studies of the sorption properties of mackinawite have yielded conflicting results, and a predictive theory for mackinawite solubility in sedimentary environments is not currently available. The presence of a solution stage in all sedimentary pyrite-forming processes (except possibly the marca- APRIL 1, 24 / ENVIRONMENTAL SCIENCE & TECHNOLOGY 135A

6 AVS measurements in sediments are not readily interpretable and have no simple meaning, although they are very easy to carry out. FIGURE site-to-pyrite transformation) has significant implications for the behavior of AVS minerals in sedimentary systems. It means that toxic metals initially adsorbed on or contained within mackinawite are released to solution during the pyrite formation process. The removal of toxic metals into pyrite then depends on the concentration of the metals in solution and their sorption characteristics relative to pyrite. The result is that, although metal toxicity can be reduced in environments in which the AVS concentration exceeds the SEM concentration, this is a temporary situation. As the metastable AVS minerals dissolve, any metals sequestered with them also dissolve and the final trace metal toxicity is determined by their behavior with respect to pyrite rather than AVS. We conclude that AVS measurements in sediments are not readily interpretable and have no simple meaning, although they are very easy to carry out. Eh ph diagram for the Fe S H 2 O system at 25 ºC The chemical dynamics of the Fe S H 2 O system determine the nature of the sulfide components of an AVS pyrite system. Pyrite nucleation in AVS systems is kinetically related to the degree of supersaturation of the solution with respect to pyrite. The graded shading of the pyrite (FeS 2 ) and FeS stability fields indicate relative supersaturation; the darker color indicates higher supersaturation. In suboxic systems, which are situated near the SO 4 2 /S( II) boundary, small changes in redox conditions (represented by the standard potential relative to the hydrogen half-cell, Eh) result in major changes in pyrite supersaturation. This means that the rate of formation of pyrite is extremely sensitive to the local physicochemical conditions. In some regions, pyrite formation will be kinetically hindered and other AVS components will dominate; in nearby regions, with slightly different redox conditions, pyrite formation may be fast and pyrite will dominate the system. (Adapted with permission from Ref. 27.) Eh (V) Fe 2+ (aq) H 2 O stability limit Fe 2+ (aq) FeS 2 FeS FeOOH SO 2 /S( II) 4 redox boundary ph Because of the potential significance of the sulfide components to the ecology of anoxic environments, further research is required on methods to more precisely determine the composition of AVS. Acknowledgments Support was provided by funds from the Louis and Elizabeth Scherck Endowed Chair in Oceanography for J. W. M and NERC grant NERLS2611 to D. R. We thank the two anonymous reviewers for their insightful comments, which led to improving this paper. John W. Morse is a professor of oceanography at Texas A&M University. David Rickard is a professor at Cardiff University in Wales. Address correspondence regarding this article to Morse at morse@ocean.tamu.edu. References (1) DiToro, D. M.; et al. Environ. Toxicol. Chem. 199, 9, (2) DiToro D. M.; et al. Environ. Sci. Technol. 1992, 26, 96. (3) Lee, B.-G.; et al. Science 2, 287, 282. (4) Lee, J.-S.; et al. Mar. Ecol. Prog. Ser. 21, 216, 129. (5) Allen R. E.; Parkes, R. J. In Geochemical Transformations of Sedimentary Sulfur; Vairamurphy, M. A., Schoonen, M. A. A., Eds.; ACS Symposium Series 612; Oxford University Press: New York, 1995; (6) Berner, R. A. Mar. Geol. 1964, 1, 117. (7) Lasorsa, B.; Caas, A. Mar. Chem. 1996, 52, 211. (8) Rickard, D.; Oldroyd, A.; Cramp. A. Estuaries 1999, 22, 693. (9) Cornwell, J. C.; Morse, J. W. Mar. Chem. 1987, 22, 193. (1) Morse, J. W.; Cornwell, J. C. Mar. Chem. 1987, 22, 55. (11) Aller, R. C. The Influence of Macrobenthos on Chemical Diagenesis of Marine Sediment. Doctoral Dissertation, Yale University, New Haven, CT, (12) Lin, S.; Morse, J. W. Am.J.Sci. 1991, 291, 55. 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