Metal Sequestration by Sedimentary Iron Sulfides

Chapter 7

Metal Sequestration by Sedimentary Iron Sulfides

Chapter Outline 3.1. The Role of the Pyrite 1. Introduction 287 Surface in Pyrite 1.1. Processes Sequestering Oxidation Trace Elements in 3.2. Metal Interactions with Sedimentary Sulfides 289 the Pyrite Surface 2. Evidence for Trace Metal 4. The Mackinawite Surface Sequestration by Sulfides in 4.1. Sequestration of Trace Sulfidic Sediments 291 Elements by Mackinawite 2.1. Trace Elements in Acid Volatile Sulfides 291 5. Sequestration of Organic Molecules by Sedimentary 2.2. Trace Elements in Pyrite 294 Sulfides 2.3. Form of Trace Element References Concentrations in Sedimentary Pyrite 297 3. The Uptake of Trace Metals by Pyrite 298

301 302 303 305

309 312

However, a satisfactory fit of the experimental points to the equation does not necessarily imply that the conditions that form the basis of the theoretical Langmuir model are fulfilled. Werner Stumm, James J. Morgan, 1970. Aquat. Chem., 454

1. INTRODUCTION The sequestration of metals in sedimentary sulfides is of considerable current scientific and technological interest. Since iron sulfides are by far the most abundant sedimentary sulfide the subject mainly involves the iron sulfides, pyrite and mackinawite. A significant amount of research has examined the role of Developments in Sedimentology. http://dx.doi.org/10.1016/B978-0-444-52989-3.00007-6 Copyright Ó 2012 Elsevier B.V. All rights reserved.

287

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the acid volatile sulfide (AVS, Chapter 14) fraction in sequestering toxic metals The sequestration of metals by mackinawite (FeSm) forms the basis of the simultaneously extracted metal (SEM)–AVS method for assessing pollution (see Fagnani et al., 2011 for a fine review in Portuguese). The SEMs are those extracted during the AVS extraction procedure. Toxicity studies suggested that when AVS > SEM, that metal is rendered nontoxic. By contrast, if SEM > AVS, the metal toxicity may be biologically significant (Di Toro et al., 1992; Di Toro et al., 1990). Although subsequent toxicological studies questioned the general applicability of this rule (e.g. Costello et al., 2011), there have been over 70 published reports using this method and it has become a standard procedure in many national environmental protection agencies. One caveat, caused by the indefinite nature of AVS components and variations in sample handling and analytical techniques (see below and Chapters 5 and 14), is that interlaboratory and interagency comparisons of SEM–AVS measurements vary by factors of several thousands (Hammerschmidt and Burton, 2010). It was thought earlier that AVS was identical to FeSm and therefore measurements of the trace element contents of FeSm (i.e. SEM) was related to the final incorporation of trace elements in pyrite in sediments. In fact, this is not the case. As noted in Chapter 14, AVS is not identical to FeSm and represents a mixture of solid and dissolved phases. In addition, because of earlier mistaken ideas that mackinawite transforms to pyrite through a solid state reaction (Chapter 6.6.1), it was assumed that the SEM metals (assumed to be sequestered in FeSm) were incorporated in resultant pyrite. Since this is not the case (i.e. because AVS s FeSm and anyway FeSm dissolves) the link between trace elements in AVS and pyrite is tenuous. The metal content of sedimentary pyrites has been of interest since not only is this a significant potential reservoir of trace metals, but it has been thought that the trace metal content of sedimentary pyrite could be used as a paleoenvironmental proxy or at least enable pyrites of hydrothermal origin to be discriminated from pyrites of normal sedimentary origin. The iron sulfides concerned in all this are pyrite and mackinawite. These are contrasting materials: pyrite is very stable with extreme insolubility whereas mackinawite is unstable and relatively soluble. Although both are sensitive to oxidation, mackinawite can be pyrophoric. This means that the surface of pyrite in sediments might prove to have regions which are not coated with oxidation products (see below), mackinawite is so sensitive to oxidation that all investigations of its surface have shown the presence of oxidation products (see Chapter 5.4.4). Finally, and even more significantly from an analytic viewpoint, sedimentary mackinawite occurs mostly as nanoparticles whereas sedimentary pyrite occurs mostly as crystals, albeit commonly microparticles. Analytically this means that pyrite surface chemistry is mainly probed by spectroscopic methods whereas these are not as widely available for mackinawite nanoparticles. Mackinawite, by contrast is relatively soluble, and its surface properties are more amenable to investigation by direct aqueous chemical techniques.

289

1. Introduction

In fact the surface properties of samples of sedimentary pyrite and sedimentary mackinawite have not been examined directly. The most direct observations are possibly the results of electron-back scatter detection spectroscopy on pyrite framboids, described in Chapter 6.7.2, which implicate surface aggregation processes for the formation of these textures, with consequences for the nature of the sedimentary pyrite surface. Additionally, many investigations of the trace element chemistry of sedimentary pyrite and mackinawite (or at least concentrates which may contain mackinawite, see Chapter 14.1.3) have been published which can indicate chemical surface properties of these materials. Most of the information regarding the surface chemistry of pyrite and mackinawite are collected with synthetic samples of mackinawite or large natural and usually hydrothermal crystals in the case of pyrite. In interpreting the results of these studies in the context of sedimentary environments, the reader should be aware that the material examined experimentally may not accurately represent sedimentary materials. For pyrite, fracture or cleavage surfaces are less likely in sediments and ideal pyrite surfaces are unlikely to be widely available; for mackinawite, the nature of the material changes with time and its sensitivity to oxidation and dissolution means that its surface behavior in sediments is variable and unlikely to be identical to that of any single precipitate.

1.1. Processes Sequestering Trace Elements in Sedimentary Sulfides This chapter was originally described as adsorption of metals on sedimentary sulfides. As is shown below, the use of the term adsorption may be misleading in this context since the reactions of metals with sedimentary sulfides are various and adsorption, sensu stricto, might be a relatively minor process. One problem has been, as suggested by the citation that heads this chapter, that results of experimental studies of metal–iron sulfide interaction can often be displayed in terms of an adsorption isotherm. However, this does not mean that the sequestration of the metal is necessarily an adsorption process. The initial reactions with the sulfide surface include physisorption, where the metal ion is weakly held to the surface with its hydration sphere (see Chapter 3.1.1) and chemisorption, where the metal is more strongly bound to the surface by exchange of water from its hydration sphere. This adsorption stage in the reaction process may be very brief and the final sequestration of the metal in sedimentary sulfides may be a consequence of a number of processes. There are five separate processes for the incorporation of trace elements in the sulfides. 1. Metal adsorption: FeSx þ M/Fe  S h M

(1)

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where FeSx refers to pyrite or mackinawite, M is a trace element and Fe–S h M is the chemiabsorbed product on the FeSx surface. 2. Trace element inclusion: FeShM/FeðMÞS þ Fe

(2)

where Fe(M)S is structurally bound M in the sulfide. The maximum possible extent of substitution of Fe or S in pyrite and mackinawite by other elements is determined experimentally by the compositions of synthetic materials usually, in the case of pyrite, at high temperatures. The data are useful in terms of defining maximum concentrations in sulfides but may have little direct relevance to interpreting the origin of compositions of sedimentary materials. It is doubtful whether the high (up to 10 wt % Cu) sometimes found in sedimentary pyrite is in solid solution. It is more probably in the form of discrete Cu sulfide minerals, as discussed below. Furthermore, improved technology has revealed that much of the trace element content of pyrite and mackinawite is in this form rather than in solid solutions. 3. Metathesis or metal exchange reactions: FeSx þ M/MS þ Fe

(3)

Exchange reactions have proven important for pyrite and mackinawite surface reactions and produce discrete sulfide phases, MS, as discussed above. 4. Coprecipitation: FeSx þ Hþ /FeðIIÞ þ SðIIÞ

(4)

M þ SðIIÞ/MS

(5)

In this sequence of reactions, sulfide produced by the dissolution of pyrite or mackinawite reacts with a metal or metalloid in solution to precipitate as a sulfide. In order for this to occur, the solubility of the new phase must be less than the solubility of mackinawite or pyrite. In fact, the difference must be considerable since the dissolved metal or metalloid must compete with the relatively high-iron concentration produced by the dissolution process for the sulfide. For this reason, the process is more prevalent with mackinawite than pyrite, since pyrite is very insoluble and the stable sulfide phase in most sedimentary environments. By contrast many metal sulfides are less soluble than mackinawite and readily form by this process. One caveat here is the kinetic factors. As noted in Chapter 6.6.5, pyrite nucleation requires extreme supersaturations (>1011 times the thermodynamic solubility product) to nucleate. This means that dissolved metals that do not require such high supersaturations can compete with pyrite. This is one explanation for the

2. Evidence for Trace Metal Sequestration by Sulfides in Sulfidic Sediments

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formation of unstable mackinawite experimentally and in some sedimentary systems, as discussed in Chapter 6.6.1. 5. Surface redox reactions: FeSx þ Mn /FeðIIIÞ þ Sð0Þ þ Mn1 M

n1

þ H2 O/MðO; OHÞ

(6) (7)

Both pyrite and mackinawite are Fe(II) minerals with reduced sulfur, mackinawite with S(II) and pyrite with more oxidized S(I), although as discussed below S(II) sites occur on the pyrite surface. Interaction of the surface with more oxidized species (such as Fe(III), Au(III), As (IV)) can result in reduction together with oxidation of the surface Fe and S species. If these more reduced species are relatively insoluble, such as Au(0), they precipitate and are incorporated in the growing minerals. The problem with nanoparticulate materials has been to distinguish sequestration by these processes from simple precipitation of distinct metal sulfide phases. It can be seen, for example, that since the solubility of some metal sulfides is less than that of mackinawite (at least in the bulk phase, see Rickard and Luther, 2006) sufficient aqueous S(II) is in equilibrium with FeSm to precipitate the relevant metal sulfides. In these nanoparticulate mixtures, it is very difficult technically to distinguish separate phases by standard methods, such as X-ray Powder Diffraction (XRPD) and chemical analysis. One way forward is to use X-ray Adsorption Spectroscopy (XAS), including Extended X-ray Absorption Fine Structure (EXAFS) and X-ray Absorption Near Edge Structure (XANES), to probe the immediate molecular environments of the metals and sulfur in the mackinawite-dominated mixture.

2. EVIDENCE FOR TRACE METAL SEQUESTRATION BY SULFIDES IN SULFIDIC SEDIMENTS 2.1. Trace Elements in Acid Volatile Sulfides SEM analyses are used together with AVS to estimate metal toxicity in the environment (e.g. Di Toro et al., 1992; Allen et al., 1993; Fagnani et al., 2011). It was found that metal toxicity in polluted aquatic sediments was better estimated by analyzing the AVS-associated metal fraction than the total metal concentration (Di Toro et al., 1992). The SEM–AVS method is an analytically simple and widely used approach to assessing potential metal toxicity in sediments. It generally consists of adding hot or cold 1 N HCl to a sediment sample and determining the concentrations of evolved hydrogen sulfide and leached metals of interest, although there are variations in the technique including acid strength, temperature, leaching time and use of antioxidants (see Chapter 14.3). The extracted fraction is equivalent to the highly reactive and

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poorly reactive iron components in the Canfield et al. (1992) protocol. The iron extracted from this technique therefore includes the Fe in sulfide complexes (e.g. FeSHþ), Fe in aqueous sulfide clusters, any mackinawite–Fe present, part of the greigite–Fe and pyrite–Fe, remnants of the nonreacted highly reactive iron, including nanoparticulate goethite and any other Fe (oxyhydr)oxides, as well as Fe in organic complexes and organic colloids, and Fe derived from sheet silicates by the activities of FeRB. As discussed in Chapters 5 and 14, the material dissolved in the AVS treatment cannot be assigned to any individual mineral, such as FeSm. The metals released by this treatment behave in a similar but not necessarily identical fashion to Fe and the exact source of these metals is element specific. The nature of the dissolved metal speciation depends partly on the sulfidation state of the system. In high sulfidation environments a greater proportion of the metal is likely to be sulfide-bound, for example. However, the age of the sample (and thus sedimentation rates) also affects the degree of sulfidation of the iron components (Raiswell and Canfield, 1996) and the behavior of the organic-Fe (Wu et al., 2001). The source of the metal constituents, which partly determine the ratios of highly reactive: poorly reactive: nonreactive metals will also contribute to the final result. The microbial ecology affects the degree to which abiologically nonreactive metal components are made available to the system, and this in turn will depend on factors such as seasonality and nutrient supply. A further factor that needs to be considered stems from the findings of Luther et al. (2001) that the formation of metal sulfide clusters helps drive the ecology of the system. Luther et al. (2001) were mainly concerned with the removal of highly toxic S(II) species from the system in the form of FeSaq clusters. However, the process also sequesters the metal component. Aqueous sulfide clusters of several other metals have been characterized (see Chapter 4.3). The effect of these aqueous metal clusters on metal toxicity estimations may be dramatic. For example, Agþ is toxic and the AVS–SEM system has been studied as a means of evaluating Ag toxicity (e.g. Berry et al., 1999). Rozan and Luther (2002) showed that the presence of aqueous Ag sulfide clusters reduced the free Agþ concentration in AVS systems from 1016 M to 1027 M. They also showed that Ag would rapidly replace Cu and Zn in aqueous Cu and Zn sulfide clusters. And, of course, AgS is relatively insoluble especially in the routine HCl treatment used in AVS analyses. The sequestration of metals in these sulfide clusters and the competitive substitution reactions between sulfide clusters, both suggest that the application of conventional SEM–AVS measurements to metal toxicity is not straightforward. The reaction might be predicted from the electrochemical constants for these soft metals. However, in systems containing a spectrum of trace metals, competition between the metals for available sulfide is expected to occur. Prediction of the effects of these potential ecological drivers in natural systems is complicated by the removal of specific components by competitive substitution reactions amongst aqueous metal sulfide clusters.

2. Evidence for Trace Metal Sequestration by Sulfides in Sulfidic Sediments

293

AVS is derived from a complex material that varies in composition depending on the environment. The sulfide that is evolved on treatment with acid probably often mainly derives from dissolved S(II) species, such as H2S and HS, Fe(II) bisulfide complexes and FeSaq complexes. The SEM concentration also varies with time and location. For those using the AVS–SEM method for assessing potential metal toxicity, this leads to the discomforting conclusion that: 1) very different dissolved and solid sulfide components may be making varying contributions to evolved H2S, and 2) that investigators using what appear to be rather minor variations in the extraction method for AVS can obtain quite different results for the same sample (e.g. hot 6 N HCl typically yields over five times the AVS of cold 1 N HCl). It is not a reasonable expectation that aqueous S(II) components and sulfide minerals undergo the same reactions with sedimentary trace metals. Therefore, it cannot be assumed that, for any given AVS-yielding S(II) concentration, the aqueous S(II) components and sulfide minerals have similar influences on the toxicity of metals (Cooper and Morse, 1998). Consequently, the “critical” AVS to SEM ratio may be quite different for sediments with different relative sources of AVS. The results of the interlaboratory comparison reported by Hammerschmidt and Burton (2010) showing extraordinary variations between reported SEM:AVS ratios, are consistent with this discussion. Because of these complex biogeochemical dynamics (Morse and Rickard, 2004), sediments can have very different AVS distribution patterns (e.g. subsurface maximum, disappearance with depth, and increase with depth) which have been observed on different length scales (see Chapter 14.1.3). This precludes there being a universally correct sampling interval for the AVS–SEM method. AVS may undergo major concentration changes on time scales of a few days to seasonally in response to hypoxia/anoxia in overlying waters. Indeed, AVS components, such as dissolved sulfide, have been shown to undergo major concentration changes in sediments on a diurnal time scale. Consequently, the AVS to SEM ratio in most coastal sediments is often likely to be constantly changing, and sampling at different times could potentially lead to quite opposite conclusions about potential metal toxicity in such sediments. Huerta-Diaz et al. (1993) reported a linear relationship between the measured AVS–Fe and the amount of trace metals associated with the AVS fraction. Oakley et al. (1980) studied the distribution of metals between various chemical phases. Copper showed a strong preference for FeS (63 %), whereas cadmium, lead and zinc were less associated with FeS (20–30 %). HamiltonTaylor et al. (1996) found that one of the most important factors controlling dissolved zinc and copper concentrations at the sediment water-interface is sulfide precipitation. Luther et al. (1980) carried out X-ray diffraction analysis of estuarine sediment samples from Newark Bay and found that the sediments contain iron monosulfides containing trace amounts of manganese, which could

294

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be due to the similar ionic radii for manganese and iron, or the precipitation of a small amount of MnS. Jacobs and Emerson (1982) found evidence for the coprecipitation of Mn with FeS and dissolved manganese is often observed in the same depth range in sediments as mackinawite formation. Much work has related cadmium toxicity to AVS (e.g. Di Toro et al., 1992; Di Toro et al., 1990). The titration of cadmium with the AVS fraction shows that cadmium can form a precipitate of cadmium sulfide. Cooper and Morse (1998) carried out acid extractions to predict the potential bioavailability of heavy metals in anoxic sediments. The bioavailability of cadmium, lead and zinc could be predicted. However predictions of nickel, copper and mercury availability could not be made, as their sulfides are not acid soluble. Typical values for AVS and SEM from sediments impacted by various forms of human activity are listed in Table 1. The data show the extreme variation in both parameters and, as discussed above, these are likely to vary seasonally and according to the sampling protocol. However, the list shows that extreme metal loads are reported from some environments, far greater than the sulfide concentrations recorded by the extraction process. General interpretations are difficult and even site specific conclusions must be treated with caution. For example, we sampled sediments in Cardiff Bay and found oregrade Zn concentrations. It turned out that the site was next to a dock which had been used for unloading bulk ore carriers when Cardiff was a major port 100 years ago. The ore-grade Zn concentrations were literally just that. More importantly, the summary in Table 1 does suggest how extensive human activities may impinge on local environments and how potentially difficult it may be to separate natural background from anthropogenic sources for metals in, especially, present-day inshore and terrestrial aquatic sediment systems.

2.2. Trace Elements in Pyrite Huerta-Diaz and Morse (1992) investigated the relationship between trace elements and pyrite in sulfidic sediments from a variety of environments. They dissolved dried sediment samples in 1 M HCl and 10 M HF to remove easily soluble minerals and silicates. The material that was left was mainly pyrite and organic matter. They then analyzed the pyrite concentrate for trace metals. As a control they compared the results of the total analysis of the total pyrite fraction and a fraction from which the residual organic matter was leached and found little difference, implying that no significant concentrations of trace metals were contained in the residual organic fraction of the sediment. The average concentrations of trace elements in pyrite from normal marine shelf and slope sediments from the Gulf of Mexico reported by Huerta-Diaz and Morse (1992) are listed in Table 2. The averages are very approximate records of the observed variations in the trace element concentrations and standard deviations on each average are mostly at least 50 % relative. Even so, the results show significant trends. Firstly, the concentrations of Cu, Ni, and Zn

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295

TABLE 1 Typical Values of AVS and SEM from Recent Sediments Variously Impacted by Human Activities AVS (mmol g1)

SEM (mmol g1)

Zinc Smelter

3.0e126

2.9e374

Belledue Harbor, Chaleur Bay, NB., Canada

Pb smelter and battery plant

5.5e102

1.9e18.4

Foundry Cove, Hudson River. NY. USA

Battery plant

0.40e64.6

0.20e779

Bear Creek, Patapsco River, MD., USA

Municipal and industrial

0.40e304

0.64e31.0

Salt Marsh, Buzzards Bay, MA., USA

Metal products manufacture

0.44e419

0.73e31.8

Steilacoom Lake, WA.,USA

Anti-algal treatment

0.02e5.65

0.60e3.91

Keweenaw Watershed, MI., USA.

Mine waste

0.006e11.6

0.36e174

Turkey creek, MO.,USA.

Mine waste

Missouri River, MO.,USA

Municipal and agricultural

2.2e20.2

0.40e2.14

Zhu river delta, NE China

Smelter

0.72e59.01

0.99e26.86

Portman Bay, SE Spain

Mine waste

47.2e510.7

16.0e144

Little Red River watershed, AR., USA

Agricultural and forest

<0.01e1.82

0.4e5.5

Begej River, Serbia

Municipal and industrial

4.3e7.1

5.2e13.8

River Danube, Serbia

Municipal and industrial

10.2e11.2

5.5e7.9

Location

Main metal source

Jinzhou Bay, Bohai Sea, China

28.1e78.2

47.6e94.5

Source: Extracted and summarized from the compilation by Fagnani et al., 2011

in pyrite in near surface samples is of the order of 1–2 %. Secondly, there is a tendency for the concentrations of all the trace elements to decrease with depth. In fact, as shown by Huerta-Diaz and Morse (1992), the concentration of the pyritic trace metals per gram dry sediment does not change significantly with depth. However, the amount of pyrite increases so that the concentration in the pyrite decreases. This is an interesting result since it shows that the more obvious cause of these trends, that the concentration of the trace element is generally more abundant in the surface sediments, is not supported by these

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TABLE 2 Average Concentrations of Selected Trace Elements (ppm) in Pyrite from Normal Marine Sediments in Surface (0e25 cm) and Depth (50e450 cm) Samples. Surface

Depth

As

2140

1106

Cd

13

7

Co

1477

588

Cr

1905

292

Cu

227,91

6869

Hg

370

52

Mn

4869

259

Mo

1181

53

Ni

9053

4130

Pb

2774

287

Zn

118,96

4058

Source: Recalculated from Huerta-Diaz and Morse (1992)

results. However, it still suggests that the earliest formed pyrite is the most trace element-rich and later formed pyrite more nearly approaches pure FeS2 in composition. This could be newly nucleated pyrite or simply the result of pyrite growth. It is unlikely to reflect trace element depletion since these do not appear to decrease with depth. The conclusion from these results is that the trace element content of sedimentary pyrites is not directly related to thermodynamic or kinetic factors but simply reflects the amount of pyrite formed and the concentration of available trace elements at any point in the sediment. So can sedimentary pyrite sequester all the available trace elements in the sediment? In order to answer this question, Huerta-Diaz and Morse (1990) introduced the concept of degree of trace element pyritization as a measure of how much of the trace element is contained in pyrite. The degree of trace element pyritization is the ratio of the trace element concentration in pyrite to the sum of the concentrations in pyrite and the HCl soluble phases. The assumption here is that the HCl-soluble phases are those that more readily supply trace elements to pore waters for incorporation in pyrite; or that the trace elements in the silicate phases are not so readily available. Their results (Huerta-Diaz and Morse, 1992) showed that the trace metals they studied could be divided into three groups whose behavior was largely independent of the

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297

sedimentary environment. Hg, Mo and As were more-or-less completely removed from the sediments by pyrite and were enriched in pyrite relative to the bulk sediment concentration. This conclusion is consistent with earlier results from Raiswell and Plant (1980) on ancient shales which showed that Hg, Mo and As were the only trace elements enriched in pyrite relative to bulk shale. Co, Cu, Mn and Ni were equally partitioned between pyrite and the reactive fraction suggesting that the concentration of these elements in pyrite reflects the chemistry of the Fe–Co–Cu–Ni–Mn–S system at low temperatures in aqueous solutions. Cr, Cd, Pb and Zn were depleted in pyrite relative to the bulk sediment suggesting that they formed stable metal–sulfide complexes in sulfidic sediments (see Chapter 4). I have spent a little time on the Huerta-Diaz and Morse (1992) results since this was the first comprehensive study of the behavior of trace metals in pyrite in recent sulfidic sediments and because their results have been largely confirmed since. However, as discussed in Chapters 13.7.8, the application of trace elements in pyrite as a paleoenvironmental indicator has not been successful to date. In the following section, I explore why this is.

2.3. Form of Trace Element Concentrations in Sedimentary Pyrite The invention of the electron microprobe in the mid-twentieth century revolutionized our understanding of the nature of trace element accumulations in sedimentary pyrite. Unfortunately, this was mainly limited to pyrite in ancient sediments. However, it did demonstrate that many of the bulk analyses of trace elements in pyrite were not as solid solutions in pyrite but in the form of discrete minerals. The electron microprobe beam diameter is of the order of 1 mm and thus nanoparticles are not discernible by this method. Since that time detailed studies of the nature of the form of trace element accumulations in pyrite have been made using high resolution techniques such as high resolution transmission electron microscopy, selected area electron diffraction, high-angle annular dark-field transmission electron microscopy and analytical electron microscopy. The nature of gold accumulations has been a target of special interest, mainly because of the economic benefits of recovering gold from pyrite concentrates. Much of the gold in pyrites is in the form of nanoparticulate native gold exsolved from pyrites on crystallization and equilibration. However, some part of the gold is also in the form of “invisible gold” (Cook and Chryssoulis, 1990), where discrete gold grains could not be identified. This invisible gold appears to be in the form of nanoparticles within the pyrite in specific regions of the pyrite crystals. Similarly, Ni and Co in pyrite is commonly concentrated in bravoitic areas in zoned pyrites. Cu is usually in the form of chalcopyrite, Zn is often observed as sphalerite particles, Pb as galena and As as arsenical pyrite or arsenopyrite.

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Deditius et al. (2011) used a battery of high-resolution methods to define the form of trace elements in pyrite from a number of low temperature ore deposits. They found that most of the pyrite trace element contents were contained in nanoparticles (<10 nm) within the pyrite. The nanoparticles included Ag–Au native metals, sulfides and sulfosalts. Deditius et al. (2011) concluded that the nanoparticles in their ore pyrites were formed primarily by exsolution from the pyrite matrix. The same chemical processes should occur in sedimentary pyrites but they have not been traced as well in these systems, mainly because of the relatively fine-grained nature of the original pyrite crystals (Chapter 6.7). In particular the form of the trace metals in the pyrite of recent sediments has not been widely identified. Even so, these studies suggest that the trace element contents of sedimentary pyrites are unlikely to be homogenously distributed. The heterogeneous distribution, in nanoparticles within the pyrite grains, means that bulkdor even microscopicdchemical analyses of trace element contents of sedimentary pyrites are likely to be highly variable. One consequence is that the use of trace element contents of pyrite as environmental proxies has not proved very fruitful to date (see Chapter 13). The results may also help explain the occurrence of Co, Ni, Cu, Zn and Pb sulfide minerals associated with pyrite in recent sediments and ancient shales (see Chapter 14). The general occurrence of trace metals in discrete minerals rather than in solid solution in sedimentary pyrite means that the use of trace elements in pyrite as paleoenvironmental proxies is likely to be unsuccessful. Apart from the problem of representative sampling of pyrites and associated minerals, the discrete mineral particles are likely to be undersaturated with respect to their dominant metals and dissolve during diagenesis. This has been observed to occur with the gold associated with deep vent pyrites. Analyses of bulk sedimentary pyrites therefore tend to give erratic and nonreproducible results.

3. THE UPTAKE OF TRACE METALS BY PYRITE The mechanisms by which pyrite initially interacts with metal ions in solution involves the pyrite surface. It has been widely studied since it is potentially important to the mobility of trace metals in sediments and may affect the biology of the system. Technological advances in the later part of the twentieth century provided a means of probing the pyrite surface at the atomic level and have lead to a deeper understanding of its roles in chemical reactions. A series of reviews have been published describing the advances that have been made as a result of these technological innovations (e.g. Rosso and Vaughan, 2006b; Rosso and Vaughan, 2006a; Murphy and Strongin, 2009). One problem with surface studies of pyrite is that pyrite has only one not very good cleavage direction along (100). The reason this presents a difficulty is that crystal cleavage is a standard means of preparing pristine surfaces for investigation. The (100) cleavage gives a conchoidal fracture surface with only limited zones

299

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of true (100) surfaces (Eggleston and Hochella, 1992). This surface is ideally charge neutral since there is a 1:1 ratio of dangling Fe to dangling S bonds and therefore relatively stable. However, surface defects give rise to monosulfide species on the pyrite (100) surface (e.g. Rosso and Vaughan, 2006a). It has been suggested that the formation of these monosulfide species is the reason behind the occurrence of Fe3þ ions on the pyrite surface (Nesbitt et al., 2000). The surface energy is a measure of the thermodynamic stability of a surface and reflects the energy required to form the surface. The lower the surface energy the more stable is the surface. Surface energies are different for different pyrite crystal faces reflecting the variable amounts of energy required to form the face and thus some faces are more stable. As shown in Table 3 the most stable face with the lowest surface energy is the cube face (100), followed by the octahedral (111) face and finally by the pyritohedral (210) face. The surface energy of pyrite is difficult to measure experimentally since most techniques (such as measuring the surface tension between water and the surface) are subject to large errors due to, especially, chemical interactions between the fluid and the surface. Thus, the surface energies for pyrite crystal faces listed in Table 3 are derived by computation using molecular modeling methods. The relevance of these measurements of ideal, perfect pyrite surfaces, to the real world are often difficult to evaluate. Pyrite crystal surfaces are not perfect and include a number of defects, such as steps and kinks, which increase the surface energy and step edges and corners may dominate surface reactions. It has become increasingly apparent that Fe and S vacancies in the surface play a significant role in determining the chemical behavior of pyrite. Fe vacancies

TABLE 3 Computed Surface Energies (g) in J m2 for Vacuum (gvac) and Hydrated (gwater) Pyrite Surfaces From Rosso and Vaughan’s (2006b) Compilation Surface

gvac

gwater

(100)

0.98e1.23

1.13

(100) crenellated

3.19

2.60

(110)

1.68e2.36

1.66

(110) facetted

1.54e1.97

1.62

(111)

1.40e1.60

e

(111) Fe

3.81

2.87

(111) S2

3.92

2.89

(210)

1.50

e

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are shown on the pyrite surface in Fig. 1. Figs 1 a and b are images taken some 60 s apart and show how dynamic the pyrite surface is: vacancies present in one image are healed in the later image by surface diffusion; atoms in the one image are not present in the other. It is attractive to surmise in Fig. 1 that Fe atom C has filled hole B, but this of course, cannot be proved. The other feature of the scanning tunneling images of the pyrite surface in Fig. 1 is the prevalence of step edges on the surface. The black areas near the figure designations a and b on the images represent lower planes of the surface which do not give a signal. The adjacent rows of white Fe atoms therefore delineate a highly reactive step edge on the surface.

FIGURE 1 Scanning tunneling microscopic image of a pyrite (100) surface under ultra high vacuum. The white areas represent Fe atoms and the bulls eye is a site registration marker. The Fe vacancy marked by B in Fig. 1a is not present 60s later in Fig 1b. By contrast the Fe atom marked by C in Fig. 1b is not present in Fig. 1a. The black areas are the result of surface steps: atoms in these areas are on a lower plane which is not picked up by the imaging tip. (Image extracted and modified from Rosso et al., 2000).

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301

3.1. The Role of the Pyrite Surface in Pyrite Oxidation Chaturvedi et al. (1996) used an in situ ultra high vacuum technique to study pyrite growth surfaces. These surfaces show a lower concentration of monosulfide species than fracture surfaces. Guevremont et al. (1997) studied the interaction of H2O, O2 and CH3OH with clean pyrite growth surfaces. They reported that H2O and CH3OH molecules are adsorbed preferentially at the small populations of defect sites on the growth surfaces, which were thought to be sulfur vacancies. The idea is that once these defect sites are saturated, the adsorbates cluster on the least reactive stoichiometric sites. Guevremont (1998) reported that O2 did not react with the pyrite surface in the absence of H2O whereas H2O oxidized the pyrite surface in the absence of O2. By contrast, fractured surfaces react with O2 (Rosso et al., 1999b). Borda et al. (2003) experimentally confirmed the original suggestion of Moses et al. (1987) that H2O2 was formed at the pyrite surface by the reaction between Fe(III) in a Sdeficient defect site (hFeIII)and adsorbed water (H2Oads) to produce an adsorbed hydroxyl radical, OH, and reduced Fe (hFeII): hFeIII þ H2 Oads /hFeII þ OH,ads þ Hþ

(8)

H2O2 is then produced by the reaction between two hydroxyl radicals: OH,ads þ OH,ads /H2 O2

(9)

H2O2 is a powerful oxidizing agent. The reaction explains many previous observations. For example, as noted above, H2O oxidizes the pyrite surface but pure, dry O2 does not. Isotopic studies show that the oxygen in the SO2 4 produced by pyrite oxidation in both abiotic and biotic systems derives from water and not O2 (Balci et al., 2007; Reedy et al., 1991; Taylor et al., 1984). Furthermore, the initial production of SO2 4 as an oxidation product of pyrite is explained by the direct reaction of OH, with nonstoichiometric S(II) at the surface. It may also explain the observation of Allen et al. (1996) that OH, occurs in acid mine drainage: OH, would be produced by the photo-dissociation of H2O2. Rosso et al. (1999a) showed that O2 oxidizes dangling Fe bond surface sites on cleaved pyrite surfaces. Eggleston et al. (1996) showed that oxidation of fractured pyrite surfaces leads to the formation of oxide patches on the surface. It appears that oxidation is initiated at specific surface sites and spreads to form randomly distributed oxidized areas next to unoxidized surface regions. The initial FeIII sites act as conduits for electron transfer from neighboring FeII to O2. Moses and Herman (1991) had shown in aqueous models of pyrite oxidation that the Fe(II)/Fe(III) couple is important for the process of electron transfer to dissolved O2. It has long been known (e.g. Singer and Stumm, 1970) that Fe(III) is a key reactant in the oxidation of pyrite in aqueous solutions. Indeed, Moses et al. (1987) showed that Fe(III) is the direct oxidant even in the presence of dissolved O2. Microorganisms, as

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discussed in Chapter 9, increase the rates of pyrite oxidation by up to 106 (e.g. Singer and Stumm, 1970) mainly by oxidizing Fe(II) to Fe(III) which then reacts as in the abiotic process. Chandra and Gerson (2011) confirmed much of the above with a detailed study of the oxidation of pyrite. They reported that both H2O and O2 are required for effective initiation of pyrite oxidation. The reaction is initiated where O2 dissociatively and H2O molecularly adsorb onto the dangling Fe bond sites on the surface. Fe oxidation occurs before S oxidation. H2O reacts to form OH, radicals which control the oxidation of S with the production of SO4(II) through thiosulfate intermediaries. They also confirmed that oxidation occurred in patches on the surface. The dominant FeIII oxide products at sedimentary pH values are iron(III) oxyhydroxides and hydroxysulfates (Todd et al., 2003) and goethite has been identified at pH 9 (England et al., 1999). It appears that the oxygen in the FeIII oxide products is mainly derived from O2 and not, as described above for SO4(II), from H2O (Usher et al., 2004). Elstinow et al. (2001) showed how important these oxide patches are to pyrite surface chemistry. They blocked FeIII sites with phosphate and showed that the rate of oxidation was significantly inhibited. It seems probable that in sedimentary environments, FeIII oxide patches on the pyrite surface are likely to have a significant effect on the reactivity of pyrite to metal species.

3.2. Metal Interactions with the Pyrite Surface In their review, Murphy and Strongin (2009) reported that studies of the adsorption and desorption of Au, Ag, Cd, Cr, Cu, Hg, Mo, Ni, Pb, Pd, Se, Sr, Eu and Zn on pyrite surfaces have been described. As noted above, there has been considerable interest in the interaction of Au with pyrite. Au(I) and Au(III) in solution are ultimately reduced at the pyrite surface to form Au0 with the oxidation of surface FeII to FeIII. Murphy and Strongin (2009) suggested that the charge transfer steps producing metallic gold would probably be located at the defect monosulfide sites or oxide patches on the pyrite surface. Redox processes at the pyrite surface dominate the reactions of many metal species with pyrite surfaces. For example, Cr(VI) is reduced to CrIII on pyrite surfaces with the production of FeIII (Mullet et al., 2007) and a Cr-rich hematite is formed. As(III) is reduced to AsII with the oxidation of FeII and SI and the formation of arsenopyrite-like structures (Bostick and Fendorf, 2003). By contrast, divalent metal species, such as Cd(II) and Hg(II) react with the pyrite surface without any redox reaction. Cd(II) reacts to form S(0), CdS and Fe(OH)2 in a disproportionation reaction (Bostick et al., 2000) and Hg(II) is absorbed on both oxidized and nonoxidized regions of the surface with the formation of Sn(II), but no Hg(0), S(II) or S(VI) species were found (Ehrhardt et al., 2000).

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The reaction of Cu(II) with pyrite has been widely studied since the products (e.g. CuS, Cu4FeS5 CuFe2S3 and CuFeS2) are dominantly CuI sulfides (Rickard and Cowper, 1994; Van der Laan et al., 1992; Luther et al., 2002). The initial adsorption reaction involves the reduction of Cu(II) to CuI (von Oertzen et al., 2007). Although these authors did not suggest this, they identified that the dominant reaction was with monosulfide defect sites of the pyrite surface and the reaction of Cu(II) with S(II) has been shown to produce Cu(I) in aqueous solutions (Luther et al., 2002). forms labile surface complexes on the pyrite surface whereas MoO2 4 forms strong inner sphere complexes with the formation of cubane-type MoS2 4 Mo–Fe–S structures analogous to the active sites of nitrogenase (Bostick et al., 2003).

4. THE MACKINAWITE SURFACE In the classic approach to a chemical description of a surface, a series of mainly acid–base titrations are made against the precipitate. It is assumed (e.g. Dzombak and Morel, 1990) that: 1. Sorption reactions at the solid interface take place at specific coordination sites. 2. Sorption reactions can be quantitatively described by conservation of mass equations. 3. Surface charge results entirely from the sorption reactions. 4. The effect of surface charge on sorption can be taken into account by applying a correction factor derived from the electric double-layer theory to constants from the conservation of mass for surface reactions. The result of these studies is a powerful model that may be used to predict the behavior of the surface. The model is quantitative and includes a series of equilibrium constants for reactions at each of the coordination sites. Wolthers et al. (2005) described a surface complexation model for nanoparticulate FeSm. They used a constant capacitance modeldwhich is a simplified form of the diffuse-layer model-to take into account of the surface charge on proton adsorption. Their experimental data showed a model with two equally concentrated surface sites: 1. ¼ FeSH0, a strongly acidic monocoordinated surface functional group which and deprotonates according to Eqns (10) and (11). ¼ FeSH0 þ Hþ 4 ¼ FeSH2 

¼ FeSH 4FeS þ H 0

þ

 log Ks1 ¼ 8:0  0:1

log Ks2

¼ 6:5  0:1

(10) (11)

where Ks1 and Ks2 are the apparent equilibrium constants for the strongly acidic site.

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2. h Fe3SH0, a weakly acidic sulfur site which protonates and deprotonates according to Eqn (12) and Eqn (13).  hFe3 SH0 þ Hþ 4hFe3 SHþ 2 log Kw1 ¼ 7:85  0:05

(12)

hFe3 SH0 4hFe3 S þ Hþ log Kw2 < 9:5

(13)

where Kw1 and Kw2 are the apparent equilibrium constants for the weakly acidic site. The surface equilibrium constants derived from this model closely describe the behavior of FeSm in aqueous solutions. As pointed out by Wolthers et al. (2005) the model is theoretical since it does not include spectroscopic evidence. Early attempts to probe the mackinawite surface with X-ray proton spectroscopy (XPS) by the Cardiff group showed that the surface of aqueous FeSm was covered in oxygen atoms and the pristine surface was not visible. This is consistent with the result of Mullet et al. (2002) who found that the surface of mackinawite contained 19 atomic % O by XPS analyses. They ascribed this to oxidation during sample handling. However, we made various attempts to eliminate O2 from both the synthesis and handling in the analytical apparatus, but to no avail. Interestingly, the surface complexation model predicts an FeSm surface with a significant oxygen concentration (Fig. 2). Naþ is the most abundant cation in marine solutions and is known to react with FeSm to form relatively stable alkali iron(II) sulfides. This would suggest a potential for Naþ to be sorbed onto FeSm surfaces. Since Naþ is far more abundant than the transition elements in most natural solutions, it may have an indirect effect on the sorption properties of FeSm relative to those elements.

FIGURE 2 Schematic representation of a mackinawite surface at the point of zero net proton charge according to the surface complexation model. The surface is viewed perpendicular to <001>. B represents S atoms with partial charges marked, l represents Fe atoms and o oxygen atoms. The sketch shows a situation where both the monocoordinated and tricoordinated surface S sites are protonated. (From Wolthers et al., 2005).

4. The Mackinawite Surface

305

Renock et al. (2009) probed the FeSm surface with a variety of techniques as part of an investigation of the sequestration of As(III) by mackinawite. They found Fe, S, O, Na and Cl on the surface. Na and Cl were residues from sample drying (cf. Rickard et al., 2006). They also noted that FeSm surfaces probably oxidized during transfer of the sample to the machine. The technical problems involved in the handling of FeSm are discussed in more detail in Chapter 5. In all XPS investigations the FeSm surface has been found to be S-enriched with Fe:S ratios varying between 1:1.3 and 1:2.1 (Herbert et al., 1998; Mullet et al., 2002; Renock et al., 2009). The variability probably reflects different syntheses and ages of the investigated materials. The excess surface sulfur is consistent with the surface complexation model (Fig. 2) although iron-loss and polysulfides have been suggested for similar sulfur enrichments on other iron sulfide surfaces (Herbert et al., 1998; Pratt et al., 1994). Transmission Mo¨ssbauer Analysis (Mullet et al., 2002) suggested the presence of Fe(III) in the FeSm structure. XPS analysis shows the presence of both FeII–S and FeIII–S species on the FeSm surface (Renock et al., 2009). As discussed in Chapter 5.4.4 with respect to greigite formation, Raman spectroscopy of the oxidation of FeSm by molecular O2 is initiated by the formation of an FeIII containing II  mackinawite, FeII13x FeIII 2x S, with the release of Fe and OH to a surface layer (Bourdoiseau et al., 2011). The problem with the spectroscopic analyses of the FeSm surface is that the results appear to depend on the way in which the samples were synthesized and how they were handled. In particular, dried mackinawite is extremely sensitive to oxidation and, even though care is taken during the synthesis, getting these samples into a machine for analysis and maintaining strictly anoxic conditions throughout is technically difficult. As noted above, we never achieved this in the Cardiff laboratory. It is possible that sedimentary FeSm is exposed to various levels of molecular O2 and that these empirical results reflect more closely the natural system. On the other hand, Mother Nature manages to maintain anoxic conditions far more readily than experimentalists can manage in our oxic atmospheres and natural mackinawite surfaces may be very little oxidized. In this case, the theoretical surface complexation model may provide a more accurate representation of reality.

4.1. Sequestration of Trace Elements by Mackinawite The observation that mackinawite derived from the high temperature monosulfide solid solution (mss) can contain significant amounts of metals other than Fe (Chapter 5.3.2) has given rise to considerable interest in FeSm as a material for sequestration of a range of often toxic elements. The original idea was that the mackinawite structure, with sheets of Fe atoms linked by S sequestered exotic ions between the sheets. I doubt this. It has proven difficult technically to synthesize low temperature FeSm with significant concentrations of exotic ions within the structure. It is more likely that the high metal concentrations in

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mackinawite derive from substitution of Fe atoms by other metals in the Fe sheets as a result of sulfide reacting at low temperatures with exsolved alloys from the cooling mss. Of course, the mackinawite structure is flexible enough to accommodate relatively high concentrations of these exotic metals. However, the observation raises the interesting question as to if sequestration of metals into FeSm can be involved in processes leading to the enhanced trace element concentrations in sediments and sedimentary rocks. The classic study of the adsorption of trace metals by FeS (presumably mainly mackinawite) by Phillips and Kraus (1965) concluded that the interactions proceeded mainly via metathesis and this has proven to be generally true (e.g. Huerta-Diaz et al., 1998). In unpublished work the Cardiff group, together with D. J. Vaughan’s group at Manchester University, investigated a number of metaldFeSm mixtures prepared by classical adsorption titrations with synchrotron techniques and the following account is colored by these results. The adsorption of Cr (VI) by FeSm has been studied by XANES spectroscopy (Boursiquot et al., 2002; Patterson et al., 1997). These studies showed that most Cr(VI) was reduced to CrIII at the FeSm surface with the formation of mixed (Fe, Cr) hydroxides and thiosulfate. XANES spectroscopy was used by Patterson et al. (1997) to determine the speciation of Cr, Fe and S on the adsorbent. Between 85 % and 100 % Cr(VI) was reduced to CrIII, with reduction occurring predominantly at the mackinawite surface. As initial Cr(VI) concentrations increased a greater percentage of total chromium was reduced. This may be due to increased oxidative dissolution of FeS with higher chromate concentrations, where the Fe(II) and S(II) ions in solution may increase the reduction of Cr(VI). The XANES sulfur spectra obtained showed that thiosulfate is the dominant sulfur phase formed in solution with small amounts of sulfate being produced. The Fe XANES spectra for reacted FeS were similar to that of g-FeOOH. Transition electron microscopy (TEM) imaging with energy dispersive spectroscopy (EDS) analysis was carried out on the solid product in order to characterize it. A 3:1 ratio of Cr:Fe in the solid product was measured and was confirmed by comparing the XANES spectra with that of a chromium standard [Cr(VI):Cr(III); 1:3]. Using these results a possible overall reaction was hypothesized: 2þ þ S2 O3 þ 6OH 3CrO2 4 þ 2FeS þ 9H2 O/4½Cr0:75 Fe0:25 ðOHÞ3 þ Fe (14) Boursiquot et al. (2002) analyzed the surface of mackinawite using XPS. A peak due to FeIII–O was produced after reaction with Cr(VI) and the FeII–S peak decreased in intensity with increasing Cr(VI) concentration, showing that FeII is oxidized to FeIII and iron (oxy) hydroxide species are formed. The Cr 2p3/2 spectrum contained three peaks which were indicative of the presence of Cr(III), thus Cr(VI) is reduced to CrIII. The S 2p spectra contained a monosulfide peak before the addition of Cr(VI). The production of additional sulfur species was observed as a result of S2 oxidation on addition of Cr(VI). The

4. The Mackinawite Surface

307

intensity of the S 2p spectra decreases with Cr(VI) concentration as sulfur species are either being removed or are being covered by an oxide layer. The O 1s spectra were made up of three species: oxide, hydroxides and water. The O 1s signal intensity increases with the initial Cr(VI) concentration. The appearance of CrIII and OH peaks and the decrease in the intensity of the S spectra indicate that it is likely that a new layer is formed on mackinawite surface consisting of Cr (oxy)hydroxide or FeIII–CrIII (oxyhydr)oxide (CrxFe1x(OH)3), in agreement with the observations of Patterson et al. (1997). In the case of Mn(II), which forms a more soluble sulfide, at least part of the Mn appears to be adsorbed on the FeSm surface with an apparent equilibrium constant of 7  104 (Morse and Arakaki, 1993). Coprecipitation experiments were carried out in a chemostat to determine the partition coefficient of Mn2þ in mackinawite. The proposed dominant reaction is: 2þ 2 þ ð1  xÞFe2þ ðaqÞ þ xMnðaqÞ þ HSðaqÞ /Feð1xÞ Mnx S þ H

(15)

The apparent partition coefficient for Mn2þ in mackinawite is 4.7  103 in low ionic strength solutions, with a reaction rate of 0.12 mmol L1 min1. The partition coefficient was independent of the mole fraction of MnS in the solid and was not affected by pH between pH 6.0 and 7.4. The influence of temperature from 5 to 25  C was investigated and the partition coefficient doubled for a 20  C increase in temperature. The partition coefficient decreased significantly with increasing precipitation rate, thus adsorption reactions may be slower than the formation of new surface sites. The ratio of adsorbed to coprecipitated Mn (9:1) remains close to constant at any reaction rate suggesting the solid must maintain a constant surface area. Our unpublished work shows that there is a switch in the process between pH 7 and 8. At pH 7, MnS does not precipitate and Mn(II) is adsorbed on the FeSm surface. At pH 8, MnS precipitates and (Mn, Fe) sulfide is produced. However, the XAS data show that the Mn environment in the coprecipitate is similar to that in MnS and the Fe remains in the same environment as mackinawite. As discussed in Chapter 4.4.2, we have found it difficult to synthesize Co sulfides in water around neutral pH at 25  C. The products were dominated by well-crystalline Co(OH)2 and it appears that this is at least partly a kinetic phenomenon since it occurred at high solution S:Co ratios. That is aqueous Co is removed rapidly as the insoluble hydroxide and the subsequent rate of Co sulfide formation is constrained by the dissolution of the hydroxide on equilibration. Some encouragement for this view came from the observation that Co sulfides were obtained under conditions where Co(OH)2 was less stable, such as acid or alkaline pH. However, the nature and consequently the stability of the Co sulfide formed at circum-neutral pH remains unknown, so that conclusions remain highly speculative about Co chemistry in sulfidic sedimentary environments. By contrast aqueous Co(II) reacts readily with FeSm at low temperatures. Our unpublished XAS studies show that the precipitate formed

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through the titration of Co(II) against FeSm is a mixture of discrete sulfide phases, including cobalt pentlandite. Huang et al. (2009) showed that the Ni sulfide formed is, nanoparticulate, nonstoichiometric and hydrated, Ni1.1S.1.5H2O. The material can be described in terms of a millerite (NiS)-like core and a hydrated Ni sulfide shell (see Chapter 4.4.3). This contrasts with FeSm which is anhydrous and stoichiometric. It also leads to further uncertainty about the nature of the first precipitated Co sulfide mentioned above. The key consequence of this report is that Ni sulfide solubility cannot be described by a simple reaction and the first formed precipitate is very different chemically from bulk nickel sulfides. This means that there is an intrinsic uncertainty in predictions of Ni sulfide behavior in aqueous sulfide systems. EXAFS analyses showed that the product of the titration of aqueous Ni2þ with FeSm produces Ni in a molecular environment similar to millerite (NiS) which is consistent with core composition of the synthesized Ni sulfide nanoparticles described above. It appears that adsorption is not an appropriate way to describe the processes involving Ni(II) and Co(II) sequestration in FeSm environments (cf. Morse and Arakaki, 1993). As suggested by the citation heading this chapter, the fit of titration results to an adsorption isotherm does not prove that the conditions of the isotherm prevail. Indeed, the fundamental conditions assumed by the Langmuir adsorption isotherm are physically highly improbable. CuS is far more insoluble than FeSm and kinetically the precipitation reaction is much faster (Rickard, 1995). Thus it is unsurprising that the titration of low concentrations of aqueous Cu(II) against FeSm results in CuS and FeS mixtures. At higher Cu concentrations a metathetic reaction occurs with the formation of chalcopyrite and other mixed metal phases such as the cubanites (Cowper and Rickard, 1988; Rickard and Cowper, 1994; Parkman et al., 1999; Cooper and Morse, 1999). It is important to note that this reaction is strictly sequential: FeS must be formed first which then reacts with dissolved Cu. If dissolved Cu and Fe mixtures are exposed simultaneously to S(II), discrete FeS and CuS products ensue; likewise if CuS is formed first and subsequently exposed to aqueous Fe(II), there is no reaction. All this is fairly consistent with the electrochemical series, the relative solubilities of FeS and Cu sulfides and the relative rates of formation of Fe and Cu sulfides. Zn also reacts with FeSm to form FeSm mixtures with wurtzite and sphalerite (Cooper and Morse, 1999) rather than surface adsorption. It appears that FeSm does not adsorb the first series transition metals, except for Mn at pH  7. Titrations of Cd(II) against FeSm showed a quantitative balance between the release of Fe(II) to solution and the removal of Cd(II) from solution, suggesting a metathetic reaction or coprecipitate with the formation of (Fe,Cd)S or discrete CdS phases(Coles et al., 2000; Parkman et al., 1999). XAS investigations suggested that the Cd phase produced was CdS (Parkman et al., 1999; Watson et al., 1995).

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Gold is adsorbed on FeSm surfaces in the form of AuIHS (Widler and Seward, 2002). XPS investigations showed that the form of gold was Au0 and that the reduction was accompanied by the formation of S(0), probably as polysulfides. Jeong et al. (2007) found that Hg(II) was adsorbed on the FeSm surface at low Hg(II) concentrations but formed discrete HgS phases at higher Hg(II) concentrations. Cooper and Morse (1999) observed both cinnabar and metacinnabar formation during the process. Jeong et al. (2010) confirmed these results using EXAFS techniques. The experimental results concur with observations of Hg association in sulfidic lake sediments (Wolfenden et al., 2005). Coles et al. (2000) showed that Pb(II) was involved in a metathetic reaction with FeSm resulting in the formation of PbS, which is consistent with what might be expected considering the relative solubilities of the bulk phases. Moyes et al. (2000) probed U adsorption on FeSm with XANES and EXAFS spectroscopy and found that uranyl was reduced at the surface producing a new U oxyhydroxide phase coupled with the oxidation of the FeSm surface. There have been a number of studies of As adsorption by FeSm, mainly because of the interest in the sequestration of toxic As in soils and sediment but also because sulfides are clearly important in As cycling in anoxic sedimentary environments (Couture et al., 2010; Harvey et al., 2002; O’Day et al., 2004; Wilkin and Ford, 2006). Bostick and Fendorf (2003) investigated arsenite uptake by the stable sulfide phases, pyrite and troilite. They suggested the formation of an arsenopyrite-like surface precipitate. Farquhar et al. (2002) found As(III) uptake on FeSm by sorption complexes and coprecipitation of As sulfides. Wolthers et al. (2005b) found that a surface complexation model described As sequestration on FeSm surfaces. Gallegaos et al. (2008, 2007) and Renock et al. (2009) suggested that As oxyanions were sorbed on FeSm and that As sulfides were coprecipitated. Wolthers et al. (2007) showed that As(V) oxidizes the FeSm–S(II) to S(0) and FeSm–Fe(II) to Fe(III). They found that As(III) inhibits crystal growth of FeSm and pyrite formation, which is consistent with As blocking sorption sites or surface defect points critical to the transformation. FeSm removes Se effectively from aqueous solution (Han et al., 2011). Han et al. (2011) found that Se(VI) was not removed as effectively as Se(IV) and that both processes could be described by adsorption isotherms. Scheinost and Chalet (2008) showed that selenite, (SeIVO32), is reduced to Fe7Se8, FeSe and Se(0) by FeSm.

5. SEQUESTRATION OF ORGANIC MOLECULES BY SEDIMENTARY SULFIDES There has been an increasing interest in the role of FeSm in sequestering organic molecules. This is partly driven by research into the early evolution of life but also appears to have significance in understanding the biogeochemistry of sedimentary sulfide systems. There have been, of course, extensive studies over

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many years on the interaction of organic molecules with pyrite and mackinawite in ores. In the case of pyrite, the purpose of the metallurgical treatment is usually to suppress the flotation of pyrite in mixed metal sulfide ores in order to increase the content of more valuable metals, such as Cu, Pb and Zn, in the concentrate. Mackinawite may be an important ore mineral in some magmatic Cu–Ni ores, such as in the Hitura deposit in Finland where it is the major Ni mineral. So the interaction of organic compounds with mackinawite has been studied again for metallurgical reasons. However, the organic chemicals are not compounds likely to be found in sediments and most of the results of these studies have little relevance to sedimentary sulfides except to confirm that organic molecules can attach to specific sites on the sulfide surfaces. The original study of relevance to sedimentary sulfides by Rickard and Butler (2001) showed that aldehydic carbonyls inhibited pyrite formation where FeSm was a reactant. Rickard and Butler showed electrochemically that the aldehydic carbonyls acted to prevent the formation of aqueous FeS clusters which are key intermediaries in pyrite formation as discussed below. It is uncertain whether this is a solution reaction or a surface reaction. Hatton and Rickard (2008) pointed out that nanoparticulate FeSm was smaller than, or about the same size as, many biochemical molecules. So the idea that these polymers attach to FeSm surfaces is misleading; rather FeSm nanoparticles attach to multiple sites on polymeric organic molecules. They showed that to FeSm bound to polymeric nucleic acids bound in the order chromosomal DNA>RNA>oligomeric DNA>deoxyadenosine monophosphate z deoxyadenosine z adenine. The linkage mechanism appears to be primarily electrostatic. Rickard et al. (2011) showed that FeSm reacted with plasmid DNA causing relaxation and denaturing of the supercoiled structure. The process, which is generally called nicking, is due to the production of sulfur free radicals. The reaction suggests that genotoxicity is a further contributory factor to the limited survival of organisms in sulfidic sedimentary environments. A number of organic molecules of potential importance to sedimentary environments interact with the pyrite surface (e.g. adenosine, 5’-AMP, phosphoglyceric acid, acetate, carbamide, ethylamine, formamide, D-ribose, adenine, alanine, cysteine and glycine) also interacted with the pyrite surface (e.g. Schoonen et al., 1999). These molecules appear to be adsorbed independently of charge and it is suggested that the reactions occurred at specific surface Fe and S sites. A summary of the results of investigations to date regarding the sequestration of elements by FeSm, is shown in Table 4. Although many elements in the periodic table have yet to be studied in this context, it appears that earlier assumptions that FeSm is an effective adsorbent sensu stricto for elements in sedimentary environments appear to be in error. FeSm appears to sequester most elements through metathetic reactions rather than adsorption. As noted above, the problem appears to have been experimental: the fit to an adsorption

5. Sequestration of Organic Molecules by Sedimentary Sulfides

TABLE 4 Summary of Processes Involved in the Sequestration of Trace Elements by FeSm in Sedimentary Environments Adsorbate

Experimental methods

Sequestration mechanism

Cr

Uptake expts XAS

Surface redox reaction:

Mn

Adsorption expts Coprecipitation expts XAS

Adsorption (pH 7). Coprecipitation (pH >7)

Co

Adsorption expts Coprecipitation expts XAS

Coprecipitation

Ni

Adsorption expts Coprecipitation expts XAS

Coprecipitation

Cu

Uptake expts. XAS, XRD

Coprecipitation

Zn

Adsorption expts XRD

Coprecipitation

As

Adsorption expts Coprecipitation expts XAS, XPS

Adsorption (pH <7) Coprecipitation (pH >7)

Se

Adsorption expts Coprecipitation expts XAS

Surface redox reaction

Tc

Coprecipitation expts XAS

Coprecipitation

Cd

Adsorption expts XRD

Coprecipitation

Re

Coprecipitation expts XAS

Coprecipitation

Au

Adsorption expts XPS

Adsorption

Hg

Adsorption expts XAS

Coprecipitation

Pb

Adsorption expts

Coprecipitation

U

Adsorption expts XAS

Adsorption (low U). Surface redox (high U)

Np

Adsorption expts XAS

Adsorption

311

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isotherm does not demonstrate that adsorption actually takes place. More recent probes of these systems, especially with synchrotron-based techniques, have shown that these elements are precipitated either as discrete phases because of the lower solubilities of their sulfides or through redox reactions with the FeSm surface. The areas where adsorption begins to be significant as a process is where the metal sulfide is more soluble than the Fe sulfide, such as low absorbate concentrations and acid pH. Organic molecules exhibit a similar range of behavior with adsorption and surface reaction being complemented by coupling of FeSm nanoparticles to large organic polymers. The lack of significant adsorption is interesting since adsorption is classically a reversible process and elements which are adsorbed may be readily eluted and released back into solution. The sequestration of elements in sedimentary systems by Fe sulfides may be more effective since it is dominated by surface redox reactions and the precipitation of discrete phases.

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